Chemistry Principles and Reactions, 7th Edition William L. Masterton Test bank
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Chapter 1--Matter and Measurements 1. An element is a substance that A. cannot be decomposed into two or more pure substances. B. is composed of two or more different substances. C. is a gas at room temperature. D. can be separated into two or more pure substances by distillation. E. can be separated into two or more substances by chromatography.
2. All of the following statements are true EXCEPT A. of the 115 known elements, 91 occur naturally. B. an element cannot be broken down into two or more pure substances. C. a compound is a substance that contains two or more elements. D. a pure compound always contains the same elements in the same mass percentages. E. another name for a heterogeneous mixture is a solution.
3. All of the following statements concerning water (H2O) are false EXCEPT A. the percentage of oxygen in H2O is independent of where the sample is obtained. B. H2O is a heterogeneous mixture. C. H2O is a homogeneous mixture. D. H2O is an element. E. H2O has properties similar to those of the elements hydrogen and oxygen.
4. Which of the following may be used to separate water into hydrogen and oxygen? A. filtration B. electrolysis C. distillation D. chromatography E. freezing
5. Which of the following is/are likely to form a homogeneous mixture? 1. 2. 3.
milk and ice cream blended together with chocolate syrup an egg combined with milk and mixed with a whisk 1 gram table salt combined with 250 mL of water
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 6. Which one of the following is most likely to be a heterogeneous mixture? A. vinegar (a mixture of acetic acid and water) B. blood C. antifreeze (a mixture of water and ethylene glycol) D. sodium chloride (table salt) dissolved in water E. the air trapped inside a car tire
7. Which term best describes liquid mercury at room temperature? A. gas B. element C. compound D. homogeneous mixture E. heterogeneous mixture
8. Which term best describes rocks, such as granite or marble? A. element B. compound C. distillation D. homogeneous mixture E. heterogeneous mixture
9. All of the following statements are false EXCEPT A. tin (Sn) is a homogeneous mixture of sulfur (S) and nitrogen (N). B. table salt is a homogeneous mixture of sodium metal and chlorine gas. C. the air trapped in a bicycle tire is a homogeneous mixture. D. sugar dissolves completely in water to give a heterogeneous mixture. E. water (H2O) is a homogeneous mixture containing hydrogen and oxygen.
10. Which of the following statements concerning water, H2O, is/are true? 1. 2. 3.
H2O is a chemical compound. H2O is a homogeneous mixture of hydrogen and oxygen. The percentage of hydrogen in H2O is independent of where the water is obtained.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 11. Sugar is a compound that dissolves readily in water. Which method is best for separating a mixture of sugar and water? A. distillation B. light absorption C. electrolysis D. gas-liquid chromatography E. filtration
12. Which method is best for removing the pulp from orange juice? A. distillation B. light absorption C. electrolysis D. gas-liquid chromatography E. filtration
13. Which one of the following substances is classified as a chemical compound? A. He B. S8 C. Na D. NO E. N2
14. Which of the following is the smallest mass? A. 1.5 ´ 108 pg B. 1.5 ´ 106 ng C. 1.5 ´ 103 mg D. 1.5 ´ 10–1 mg E. 1.5 ´ 10–5 g
15. Which of the following is the largest volume? A. 5.0 ´ 102 cm3 B. 5.0 ´ 10-4 L C. 5.0 ´ 103 mL D. 5.0 ´ 10-1 L E. 5.0 ´ 105 mL
16. The radius of a potassium atom is 0.227 nm. What is the radius in millimeters? A. 2.27 ´ 10–9 mm B. 2.27 ´ 10–8 mm C. 2.27 ´ 10–7 mm D. 2.27 ´ 10–6 mm E. 2.27 ´ 10-4 mm
17. Which method is correct for converting kelvin to Celsius?
A.
B. C. D. E.
18. If the temperature of a beaker of water is 65.0F, what is its temperature in Celsius? (Remember: water melts at 0C and 32F; water boils at 100C and 212F) A. 4.11C B. 18.3C C. 36.1C D. 59.4C E. 68.1C
19. If the outdoor temperature is 17.0C, what is the temperature in Fahrenheit? (Remember: water melts at 0C and 32F; water boils at 100C and 212F) A. -1.40F B. 30.6F C. 41.4F D. 62.6F E. 74.6F
20. The boiling point of ammonia is -33.3C. What is this temperature in kelvin? A. -306.5 K B. -33.3 K C. 239.9 K D. 306.5 K E. 331.5 K
21. Water freezes at 0.0C. What temperature does this correspond to in kelvin? A. 173.2 K B. 200.2 K C. 273.2 K D. 300.2 K E. 373.2 K
22. Many experiments are conducted at 298 K. What is this temperature in Celsius? A. 0C B. 25C C. 55C D. 273C E. 298C
23. The boiling point for liquid oxygen is 90.0 K. What is the boiling point in Fahrenheit? A. -361.4F B. -329.4F C. -297.8F D. -183.0F E. -151.0F
24. Express 0.00720 in exponential notation. A. 7.20 ´ 103 B. 7.2 ´ 10-3 C. 7.20 ´ 10–3 D. 7.2 ´ 103 E. 7.2000 ´ 10-3
25. Convert 8.900´ 10–8 meters to nanometers and express the answer in standard notation using the correct number of significant figures. A. 89 nm B. 89.0 nm C. 89.00 nm D. 8.90 nm E. 0.8900 nm
26. How many significant figures are in the following mass: 0.00047800 kg? A. 3 B. 5 C. 6 D. 8 E. 9
27. How many significant figures are in the following volume: 5.00 ´ 104 mL? A. 1 B. 2 C. 3 D. 4 E. 5
28. What is the correct answer to the following expression: (205.18 – 197.3) ¸ 6.226? Carry out the subtraction operation first. A. 1 B. 1.3 C. 1.27 D. 1.266 E. 1.2657
29. What is the correct answer to the expression below? 1.472 ´ 10-7 + 4.32 ´ 10-9 = A. 2 ´ 10-7 B. 1.5 ´ 10-7 C. 1.52 ´ 10-7 D. 1.515 ´ 10-7 E. 1.5152 ´ 10-7
30. What is the correct answer to the following expression? 7.4576 ´ 10–2 + 4.11 ´ 10–5 + 6 ´ 10–4 = A. 7. ´ 10–2 B. 7.5 ´ 10–2 C. 7.52 ´ 10–2 D. 7.522 ´ 10–2 E. 7.5217 ´ 10–2
31. Round the answer to the following problem to the correct number of significant figures. (14.0186 ´ 0.00458) + (15.0032 ´ 0.99542) = 14.99869 A. 15.0 B. 15.00 C. 14.999 D. 14.9987 E. 14.99869
32. A standard sheet of paper is 8.5 ´ 11 inches. What is the surface area, in cm2, of one side of a sheet of paper? (2.54 cm = 1.00 inch) A. 14 cm2 B. 37 cm2 C. 94 cm2 D. 240 cm2 E. 6.0 ´ 102 cm2
33. The dimensions of a box are 1.2 feet by 0.50 feet by 0.75 feet. Calculate the volume of the box in cubic centimeters. (2.54 cm = 1.00 inch, 12.0 inches = 1.00 foot) A. 14 cm3 B. 47 cm3 C. 306 cm3 D. 418 cm3 E. 1.3 ´ 104 cm3
34. How many miles are covered in a 15 km race? (1 mile = 5280 feet, 12 inches = 1 foot, 1 inch = 2.54 cm) A. 7.1 mile B. 9.3 mile C. 11 mile D. 15 mile E. 26 mile
35. Atomic dimensions are often reported in Ångstroms (1 Å = 1 ´ 10-10 m). If the atomic radius of an aluminum atom is 1.43 Ångstroms, what is its radius in millimeters? A. 1.43 ´ 10-13 mm B. 1.43 ´ 10-7 mm C. 0.143 mm D. 14.3 mm E. 143 mm
36. If the fuel efficiency of an automobile is 32 miles per gallon, what is its fuel efficiency in kilometers per liter? (1 km = 0.621 mile, 1.000 L = 1.057 quarts, 4 quarts = 1 gallon) A. 5.3 km/L B. 14 km/L C. 20 km/L D. 75 km/L E. 2.0 ´ 102 km/L
37. The volume of a carbon atom is 1.9 ´ 10–30 m3. What is the radius of the atom in picometers? The volume of a sphere is (4/3)pr3. A. 77 pm B. 520 pm C. 770 pm D. 3.0 ´ 102 pm E. 52 pm
38. Which of the following observations is/are examples of chemical change? 1. 2. 3.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 2 and 3
Iron (Fe) rusts, forming Fe2O3. The density of water increases when it changes from a solid to a liquid. Sodium chloride melts at 801C.
39. All of the following are examples of chemical change EXCEPT A. the fermentation of wine. B. the tarnishing of silver. C. the condensation of steam. D. the combustion of butane gas. E. the rusting of iron.
40. Which one of the following statements is not a comparison of physical properties? A. Mercury and gallium are both liquids at 50C. B. Oxygen is more soluble in water than helium. C. The density of gold is greater than the density of silver. D. Oxygen and nitrogen are both liquids at -200C. E. Calcium dissolves more quickly than iron in acids.
41. All of the following are examples of an intensive property EXCEPT A. solubility. B. boiling point. C. electrical conductivity of an element. D. density. E. both boiling point and density.
42. An intensive property of a substance is A. independent of the amount present. B. dependent on its volume, but not its mass. C. not affected by its temperature. D. dependent only on its temperature. E. dependent only on its mass and volume.
43. An extensive property is A. used to identify substances. B. independent of amount. C. related to density. D. dependent upon amount. E. observed throughout a substance.
44. 525 mL of water at 25C (density = 0.997 g/mL) is placed in a container. The water is then cooled to form ice at -10C (density = 0.917 g/mL). What is the mass and volume of the ice? A. 523 g and 525 mL B. 523 g and 571 mL C. 527 g and 525 mL D. 527 g and 571 mL E. not enough information given to solve the problem
45. If the density of nitrogen in air is 0.87 g/L, what mass (in kg) of nitrogen is contained in a room with dimensions of 4.0 m ´ 3.5 m ´ 2.4 m? A. 29 kg B. 39 kg C. 2.6 ´ 10-2 kg D. 2.9 ´ 104 kg E. 26 kg
46. A solid with a mass of 19.3 g is added to a graduated cylinder filled with water to the 25.0 mL mark. After the solid sinks to the bottom, the water level is at 35.8 mL. What is the density of the solid? A. 0.539 g/mL B. 0.560 g/mL C. 1.79 g/mL D. 2.19 g/mL E. 8.50 g/mL
47. A barometer is filled with a cylindrical column of mercury that is 76.0 cm high and 1.000 cm in diameter. If the density of mercury is 13.53 g/cm3, what is the mass of mercury in the column? A. 0.227 g B. 4.41 g C. 808 g D. 1.03 ´ 103 g E. 3.23 ´ 103 g
48. Calcium carbonate, or limestone, is relatively insoluble in water. At 25C, only 5.8 mg will dissolve in 1.0 liter of water. What volume of water is needed to dissolve 5.0 g of calcium carbonate? A. 4.6 ´ 10–3 L B. 3.0 ´ 10–2 L C. 1.4 ´ 102 L D. 3.4 ´ 102 L E. 8.6 ´ 102 L
49. At 0C, 35.7 g of sodium chloride (NaCl) will dissolve in 1.00 ´ 102 mL of water. What mass of sodium chloride will dissolve in 7.75 L of water (at 0C)? A. 4.61 g B. 4.61 ´ 101 g C. 7.75 ´ 101 g D. 2.17 ´ 102 g E. 2.77 ´ 103 g
50. At what point is the temperature in C twice that of the temperature in F? A. -40.0C B. -32.4C C. -24.6C D. -16.2C E. -8.88C
Chapter 1--Matter and Measurements Key
1. An element is a substance that A. cannot be decomposed into two or more pure substances. B. is composed of two or more different substances. C. is a gas at room temperature. D. can be separated into two or more pure substances by distillation. E. can be separated into two or more substances by chromatography.
2. All of the following statements are true EXCEPT A. of the 115 known elements, 91 occur naturally. B. an element cannot be broken down into two or more pure substances. C. a compound is a substance that contains two or more elements. D. a pure compound always contains the same elements in the same mass percentages. E. another name for a heterogeneous mixture is a solution.
3. All of the following statements concerning water (H2O) are false EXCEPT A. the percentage of oxygen in H2O is independent of where the sample is obtained. B. H2O is a heterogeneous mixture. C. H2O is a homogeneous mixture. D. H2O is an element. E. H2O has properties similar to those of the elements hydrogen and oxygen.
4. Which of the following may be used to separate water into hydrogen and oxygen? A. filtration B. electrolysis C. distillation D. chromatography E. freezing
5. Which of the following is/are likely to form a homogeneous mixture? 1. 2. 3.
milk and ice cream blended together with chocolate syrup an egg combined with milk and mixed with a whisk 1 gram table salt combined with 250 mL of water
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 6. Which one of the following is most likely to be a heterogeneous mixture? A. vinegar (a mixture of acetic acid and water) B. blood C. antifreeze (a mixture of water and ethylene glycol) D. sodium chloride (table salt) dissolved in water E. the air trapped inside a car tire
7. Which term best describes liquid mercury at room temperature? A. gas B. element C. compound D. homogeneous mixture E. heterogeneous mixture
8. Which term best describes rocks, such as granite or marble? A. element B. compound C. distillation D. homogeneous mixture E. heterogeneous mixture
9. All of the following statements are false EXCEPT A. tin (Sn) is a homogeneous mixture of sulfur (S) and nitrogen (N). B. table salt is a homogeneous mixture of sodium metal and chlorine gas. C. the air trapped in a bicycle tire is a homogeneous mixture. D. sugar dissolves completely in water to give a heterogeneous mixture. E. water (H2O) is a homogeneous mixture containing hydrogen and oxygen.
10. Which of the following statements concerning water, H2O, is/are true? 1. 2. 3.
H2O is a chemical compound. H2O is a homogeneous mixture of hydrogen and oxygen. The percentage of hydrogen in H2O is independent of where the water is obtained.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 11. Sugar is a compound that dissolves readily in water. Which method is best for separating a mixture of sugar and water? A. distillation B. light absorption C. electrolysis D. gas-liquid chromatography E. filtration
12. Which method is best for removing the pulp from orange juice? A. distillation B. light absorption C. electrolysis D. gas-liquid chromatography E. filtration
13. Which one of the following substances is classified as a chemical compound? A. He B. S8 C. Na D. NO E. N2
14. Which of the following is the smallest mass? A. 1.5 ´ 108 pg B. 1.5 ´ 106 ng C. 1.5 ´ 103 mg D. 1.5 ´ 10–1 mg E. 1.5 ´ 10–5 g
15. Which of the following is the largest volume? A. 5.0 ´ 102 cm3 B. 5.0 ´ 10-4 L C. 5.0 ´ 103 mL D. 5.0 ´ 10-1 L E. 5.0 ´ 105 mL
16. The radius of a potassium atom is 0.227 nm. What is the radius in millimeters? A. 2.27 ´ 10–9 mm B. 2.27 ´ 10–8 mm C. 2.27 ´ 10–7 mm D. 2.27 ´ 10–6 mm E. 2.27 ´ 10-4 mm
17. Which method is correct for converting kelvin to Celsius?
A.
B. C. D. E.
18. If the temperature of a beaker of water is 65.0F, what is its temperature in Celsius? (Remember: water melts at 0C and 32F; water boils at 100C and 212F) A. 4.11C B. 18.3C C. 36.1C D. 59.4C E. 68.1C
19. If the outdoor temperature is 17.0C, what is the temperature in Fahrenheit? (Remember: water melts at 0C and 32F; water boils at 100C and 212F) A. -1.40F B. 30.6F C. 41.4F D. 62.6F E. 74.6F
20. The boiling point of ammonia is -33.3C. What is this temperature in kelvin? A. -306.5 K B. -33.3 K C. 239.9 K D. 306.5 K E. 331.5 K
21. Water freezes at 0.0C. What temperature does this correspond to in kelvin? A. 173.2 K B. 200.2 K C. 273.2 K D. 300.2 K E. 373.2 K
22. Many experiments are conducted at 298 K. What is this temperature in Celsius? A. 0C B. 25C C. 55C D. 273C E. 298C
23. The boiling point for liquid oxygen is 90.0 K. What is the boiling point in Fahrenheit? A. -361.4F B. -329.4F C. -297.8F D. -183.0F E. -151.0F
24. Express 0.00720 in exponential notation. A. 7.20 ´ 103 B. 7.2 ´ 10-3 C. 7.20 ´ 10–3 D. 7.2 ´ 103 E. 7.2000 ´ 10-3
25. Convert 8.900´ 10–8 meters to nanometers and express the answer in standard notation using the correct number of significant figures. A. 89 nm B. 89.0 nm C. 89.00 nm D. 8.90 nm E. 0.8900 nm
26. How many significant figures are in the following mass: 0.00047800 kg? A. 3 B. 5 C. 6 D. 8 E. 9
27. How many significant figures are in the following volume: 5.00 ´ 104 mL? A. 1 B. 2 C. 3 D. 4 E. 5
28. What is the correct answer to the following expression: (205.18 – 197.3) ¸ 6.226? Carry out the subtraction operation first. A. 1 B. 1.3 C. 1.27 D. 1.266 E. 1.2657
29. What is the correct answer to the expression below? 1.472 ´ 10-7 + 4.32 ´ 10-9 = A. 2 ´ 10-7 B. 1.5 ´ 10-7 C. 1.52 ´ 10-7 D. 1.515 ´ 10-7 E. 1.5152 ´ 10-7
30. What is the correct answer to the following expression? 7.4576 ´ 10–2 + 4.11 ´ 10–5 + 6 ´ 10–4 = A. 7. ´ 10–2 B. 7.5 ´ 10–2 C. 7.52 ´ 10–2 D. 7.522 ´ 10–2 E. 7.5217 ´ 10–2
31. Round the answer to the following problem to the correct number of significant figures. (14.0186 ´ 0.00458) + (15.0032 ´ 0.99542) = 14.99869 A. 15.0 B. 15.00 C. 14.999 D. 14.9987 E. 14.99869
32. A standard sheet of paper is 8.5 ´ 11 inches. What is the surface area, in cm2, of one side of a sheet of paper? (2.54 cm = 1.00 inch) A. 14 cm2 B. 37 cm2 C. 94 cm2 D. 240 cm2 E. 6.0 ´ 102 cm2
33. The dimensions of a box are 1.2 feet by 0.50 feet by 0.75 feet. Calculate the volume of the box in cubic centimeters. (2.54 cm = 1.00 inch, 12.0 inches = 1.00 foot) A. 14 cm3 B. 47 cm3 C. 306 cm3 D. 418 cm3 E. 1.3 ´ 104 cm3
34. How many miles are covered in a 15 km race? (1 mile = 5280 feet, 12 inches = 1 foot, 1 inch = 2.54 cm) A. 7.1 mile B. 9.3 mile C. 11 mile D. 15 mile E. 26 mile
35. Atomic dimensions are often reported in Ångstroms (1 Å = 1 ´ 10-10 m). If the atomic radius of an aluminum atom is 1.43 Ångstroms, what is its radius in millimeters? A. 1.43 ´ 10-13 mm B. 1.43 ´ 10-7 mm C. 0.143 mm D. 14.3 mm E. 143 mm
36. If the fuel efficiency of an automobile is 32 miles per gallon, what is its fuel efficiency in kilometers per liter? (1 km = 0.621 mile, 1.000 L = 1.057 quarts, 4 quarts = 1 gallon) A. 5.3 km/L B. 14 km/L C. 20 km/L D. 75 km/L E. 2.0 ´ 102 km/L
37. The volume of a carbon atom is 1.9 ´ 10–30 m3. What is the radius of the atom in picometers? The volume of a sphere is (4/3)pr3. A. 77 pm B. 520 pm C. 770 pm D. 3.0 ´ 102 pm E. 52 pm
38. Which of the following observations is/are examples of chemical change? 1. 2. 3.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 2 and 3
Iron (Fe) rusts, forming Fe2O3. The density of water increases when it changes from a solid to a liquid. Sodium chloride melts at 801C.
39. All of the following are examples of chemical change EXCEPT A. the fermentation of wine. B. the tarnishing of silver. C. the condensation of steam. D. the combustion of butane gas. E. the rusting of iron.
40. Which one of the following statements is not a comparison of physical properties? A. Mercury and gallium are both liquids at 50C. B. Oxygen is more soluble in water than helium. C. The density of gold is greater than the density of silver. D. Oxygen and nitrogen are both liquids at -200C. E. Calcium dissolves more quickly than iron in acids.
41. All of the following are examples of an intensive property EXCEPT A. solubility. B. boiling point. C. electrical conductivity of an element. D. density. E. both boiling point and density.
42. An intensive property of a substance is A. independent of the amount present. B. dependent on its volume, but not its mass. C. not affected by its temperature. D. dependent only on its temperature. E. dependent only on its mass and volume.
43. An extensive property is A. used to identify substances. B. independent of amount. C. related to density. D. dependent upon amount. E. observed throughout a substance.
44. 525 mL of water at 25C (density = 0.997 g/mL) is placed in a container. The water is then cooled to form ice at -10C (density = 0.917 g/mL). What is the mass and volume of the ice? A. 523 g and 525 mL B. 523 g and 571 mL C. 527 g and 525 mL D. 527 g and 571 mL E. not enough information given to solve the problem
45. If the density of nitrogen in air is 0.87 g/L, what mass (in kg) of nitrogen is contained in a room with dimensions of 4.0 m ´ 3.5 m ´ 2.4 m? A. 29 kg B. 39 kg C. 2.6 ´ 10-2 kg D. 2.9 ´ 104 kg E. 26 kg
46. A solid with a mass of 19.3 g is added to a graduated cylinder filled with water to the 25.0 mL mark. After the solid sinks to the bottom, the water level is at 35.8 mL. What is the density of the solid? A. 0.539 g/mL B. 0.560 g/mL C. 1.79 g/mL D. 2.19 g/mL E. 8.50 g/mL
47. A barometer is filled with a cylindrical column of mercury that is 76.0 cm high and 1.000 cm in diameter. If the density of mercury is 13.53 g/cm3, what is the mass of mercury in the column? A. 0.227 g B. 4.41 g C. 808 g D. 1.03 ´ 103 g E. 3.23 ´ 103 g
48. Calcium carbonate, or limestone, is relatively insoluble in water. At 25C, only 5.8 mg will dissolve in 1.0 liter of water. What volume of water is needed to dissolve 5.0 g of calcium carbonate? A. 4.6 ´ 10–3 L B. 3.0 ´ 10–2 L C. 1.4 ´ 102 L D. 3.4 ´ 102 L E. 8.6 ´ 102 L
49. At 0C, 35.7 g of sodium chloride (NaCl) will dissolve in 1.00 ´ 102 mL of water. What mass of sodium chloride will dissolve in 7.75 L of water (at 0C)? A. 4.61 g B. 4.61 ´ 101 g C. 7.75 ´ 101 g D. 2.17 ´ 102 g E. 2.77 ´ 103 g
50. At what point is the temperature in C twice that of the temperature in F? A. -40.0C B. -32.4C C. -24.6C D. -16.2C E. -8.88C
Chapter 2--Atoms, Molecules, and Ions 1. Which of the following statements is/are postulates of Dalton’s atomic theory? 1. 2. 3.
Elements are composed of tiny particles called atoms. No atom is changed into another element in an ordinary chemical reaction. Compounds are formed when two or more atoms combine.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2 and 3 2. J.J. Thomson determined that electrons are small, negatively charged particles by A. bombarding gold foil with alpha particles. B. exposing photographic plates to radioactive uranium. C. deflecting cathode rays with electric and magnetic fields. D. converting cathode rays to electron particles using a fluorescent screen. E. decomposing neutrons into protons and electrons.
3. Which of the following statements is/are CORRECT? 1. 2. 3.
A neutron is an uncharged particle with a mass slightly greater than a proton. The nucleus of an atom has a negative charge. More than 99.9% of an atom's mass is concentrated in the electron cloud surrounding the nucleus.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2 and 3 4. Rank the subatomic particles from least to greatest mass. A. electrons = neutrons = protons B. electrons = protons < neutrons C. electrons < neutrons = protons D. electrons < protons < neutrons E. electrons < neutrons < protons
5. All of the following statements are true EXCEPT A. all atoms of a given element have the same mass number. B. for any neutral element, the number of electrons is equal to the number of protons. C. the mass number is the sum of the number of protons and neutrons. D. isotopes of atoms contain the same number of protons but a different number of neutrons. E. the atomic number equals the number of protons in an atom.
6. All atoms of the same element have the same number of ____. A. neutrons B. protons C. protons and neutrons D. electrons and neutrons E. protons, neutrons, and electrons
7. What is the mass number of an argon atom with 22 neutrons? A. 2 B. 18 C. 22 D. 40 E. 39.95
8. Which of the following atoms contains the fewest protons? A. 232Th B. 231Pa C. 245Pu D. 238U E. 232Pa
9. How many protons, neutrons, and electrons are in a silver atom with a mass number of 108? A. 47 protons, 47 neutrons, 61 electrons B. 47 protons, 61 neutrons, 47 electrons C. 61 protons, 47 neutrons, 47 electrons D. 47 protons, 108 neutrons, 47 electrons E. 61 protons, 108 neutrons, 61 electrons
10. F-20, a radioactive isotope of fluorine, has A. 9 protons, 10 neutrons, and 1 electron. B. 9 protons, 10 neutrons, and 9 electrons. C. 9 protons, 11 neutrons, and 9 electrons. D. 10 protons, 9 neutrons, and 1 electron. E. 10 protons, 10 neutrons, and 10 electrons.
11. Which two atoms below have the same number of neutrons? , A.
and
B.
and
C.
and
D.
and
E.
and
,
,
,
12. What is the atomic symbol for an element with 39 protons and 50 neutrons? A. B. C. D. E.
13. Which two of the ions below have the same number of electrons? ,
,
A.
and
B.
and
C.
and
D.
and
E.
and
,
,
14. Two isotopes of a given element will have the same number of ____, but a different number of ____ in their nucleus. A. protons, electrons B. electrons, protons C. protons, neutrons D. neutrons, protons E. electrons, neutrons
15. Which of the following are a pair of isotopes? A.
and
B.
and
C.
and
D.
and
E.
and
16. Which species has 63 neutrons? A. B. C. D. E. none of the above
17. Two isotopes of chlorine are found in nature, Cl-35 and Cl-37. The average mass of chlorine is 35.45 amu. The more abundant isotope of Cl has A. 17 protons, 17 electrons, and 18 neutrons. B. 17 protons, 17 electrons, and 18.45 neutrons. C. 17 protons, 17 electrons, and 20 neutrons. D. 18 protons, 18 electrons, and 17 neutrons. E. 19 protons, 19 electrons, and 16 neutrons.
18. What is the identity of A. zinc B. silver C. iridium D. cesium E. manganese
?
19. What is the symbol for an element which contains 57 neutrons and has a mass number of 101? A. Er B. Ru C. Md D. La E. Os
20. Rubidium has two naturally occurring isotopes. The average mass of Rb is 85.4678 amu. If 72.15% of Rb is found as Rb-85 (84.9117 amu), what is the mass of the other isotope? A. 0.56 amu B. 85.68 amu C. 86.91 amu D. 86.02 amu E. 83.47 amu
21. Silver has two stable isotopes with masses of 106.90509 amu and 108.9047 amu. The average molar mass of silver is 107.868 amu. What is the percent abundance of each isotope? A. 50.0% Ag-107 and 50.0% Ag-109 B. 51.8% Ag-107 and 48.2% Ag-109 C. 55.4% Ag-107 and 44.6% Ag-109 D. 48.2% Ag-107 and 51.8% Ag-109 E. 44.6% Ag-107 and 55.4% Ag-109
22. Gallium has an average atomic mass of 69.7 amu. In a typical sample, 60.4% of Ga exists as Ga-69 (68.9257 amu). What is the identity and the atomic mass of the other isotope? A.
; 70.9 amu
B.
; 70.9 amu
C.
; 70.9 amu
D.
; 71.9 amu
E.
; 71.9 amu
23. An element has three naturally occurring isotopes with the following abundances and masses:
abundance 78.99% 10.00% 11.01%
Determine the molar mass of the element.
A. 24.31 g/mol B. 24.98 g/mol C. 74.95 g/mol D. 2431 g/mol E. none of the above
mass (amu) 23.985042 24.985837 25.982593
24. The average molar mass of lithium is 6.941. A sample of lithium consists of two isotopes with masses of 6.01512 amu and 7.01600 amu. Determine the percent abundance of each isotope. A. 7.49% Li-6 and 92.51% Li-7 B. 8.45% Li-6 and 91.55% Li-7 C. 12.49% Li-6 and 87.51% Li-7 D. 91.55% Li-6 and 8.45% Li-7 E. 92.51% Li-6 and 7.49% Li-7
25. What is the mass (in grams) of a boron atom? A. 10.8 g B. 1.80 ´ 10-23 g C. 1.66 ´ 10-24 g D. 5.57 ´ 1022 g E. 1.54 ´ 10-25 g
26. Group 1 elements are also known as A. alkaline earth metals. B. alkali metals. C. chalcogens. D. halogens. E. noble gases.
27. How many nonmetallic elements are there in group 13? A. 0 B. 1 C. 2 D. 3 E. 4
28. Identify the halogen from period 4. A. Br B. I C. Kr D. Ar E. K
29. What element is in the fourth period in Group 3A? A. Sb B. Ga C. In D. Si E. Tl
30. Which group of three elements contains a nonmetal, a metal, and a metalloid? A. Li, Al, Si B. Na, Hg, I C. I, Hg, Si D. K, O, Br E. H, Al, N
31. Which three elements are likely to have similar chemical and physical properties? A. boron, silicon, and germanium B. sodium, magnesium, and aluminum C. sodium, potassium, and rubidium D. oxygen, sulfur, and chlorine E. carbon, nitrogen, and oxygen
32. Which group of three elements contains a transition metal, a halogen, and a noble gas? A. S, I, Cu B. Br, Kr, Ba C. Ar, Hg, Rn D. Ce, N, He E. Cu, I, Xe
33. How many elements are contained in period 2? A. 3 B. 8 C. 10 D. 18 E. 32
34. Of the naturally occurring elements in group 14, how many are nonmetals, metalloids, and metals? A. 0 nonmetals, 3 metalloids, and 2 metals B. 1 nonmetal, 2 metalloids, and 2 metals C. 2 nonmetals, 2 metalloids, and 1 metal D. 2 nonmetals, 1 metalloid, and 2 metals E. 3 nonmetals, 0 metalloids, and 2 metals
35. Which two of the following elements are abundant in the Earth's crust, but missing from the human body: O, Al, Si, Fe, C, N? A. O and Fe B. Si and C C. Al and Si D. O and N E. Fe and N
36. The formula of ethanol, CH3CH2OH, is an example of a(n) A. condensed formula. B. empirical formula. C. structural formula. D. ionic compound formula. E. mass spectrum.
37. Which element is most likely to form an ion with a -2 charge? A. K B. Mg C. P D. Br E. S
38. Which atom is most likely to form an ion with a +2 charge? A. scandium B. calcium C. aluminum D. oxygen E. fluorine
39. A strontium ion has ____ electrons. A. 35 B. 36 C. 37 D. 38 E. 39
40. For a nonmetal in Group 16 of the periodic table, the most common monatomic ion will have a charge of ____. A. –3 B. –2 C. –1 D. +1 E. +2
41. Identify the ions and their charges in Na2SO4. A. Na+, SO4B. Na+, SO42C. Na+, SO4D. Na2+, SO4E. Na2+, SO42-
42. Identify the ions and their charges in KH2PO4. A. K+, H+, P3-, O2B. K+, H2+, P3-, O8C. K+, H22+, P-1, O4-2 D. K+, H2PO4E. K+, H2+, PO43-
43. What is the correct formula for an ionic compound that contains magnesium ions and phosphide ions? A. MgP B. MgP2 C. Mg3P2 D. Mg3(PO4)2 E. Mg2P3
44. What is the correct formula for an ionic compound that contains aluminum ions and chloride ions? A. AlCl B. AlCl2 C. AlCl3 D. Al2Cl3 E. Al3Cl2
45. What are the values for x and y, respectively, in CaxHyPO4? A. 1 and 2 B. 2 and 1 C. 1 and 3 D. 2 and 2 E. 1 and 1
46. Sodium sulfate has the chemical formula Na2SO4. Based on this information, the formula for chromium(III) sulfate is ____. A. CrSO4 B. Cr(SO4)3 C. Cr2(SO4)3 D. Cr2SO4 E. Cr3(SO4)2
47. What is the correct name for MnS? A. manganese sulfide B. dimanganese sulfate C. dimanganese sulfide D. manganese(II) sulfate E. manganese(II) sulfide
48. What is the correct name for K3PO4? A. tripotassium phosphate B. potassium(I) monophosphorus tetraoxide C. potassium(I) phosphate D. potassium phosphate E. potassium phosphide
49. What is the correct name for TiCl4? A. monotitanium tetrachloride B. tetrachlorine titanate C. titanium tetrachlorine D. titanium(IV) tetrachloride E. titanium(IV) chloride
50. What is the correct name for Al2O3? A. alum B. aluminum trioxide C. aluminum ozinide D. aluminum oxide E. dialuminum trioxide
51. What is the correct formula for aluminum selenide? A. AlSe B. AlSe2 C. Al2Se D. Al2Se3 E. Al3Se2
52. What is the correct formula for iron(II) nitrate? A. Fe2(NO3)2 B. Fe2NO3 C. Fe(NO3)2 D. Fe3N2 E. FeNO3
53. What is the correct formula for barium perchlorate? A. BaClO4 B. BaClO3 C. Ba(ClO4)2 D. Ba(ClO3)2 E. Ba(ClO3)3
54. What is the correct name for N2O3? A. nitrogen oxide B. nitrogen(II) oxide C. nitrogen(III) oxide D. trioxygen dinitride E. dinitrogen trioxide
55. What is the correct name for PF5? A. phosphorus pentafluoride B. phosphorus(V) fluoride C. phosphorofluoride D. pentafluorophosphorus E. pentafluorophosphate
56. What is the correct name for CCl4? A. carbon chlorine B. tetracarbon chloride C. carbon tetrachloride D. carbon(IV) chloride E. tetrachlorocarbide
57. What is the correct formula for sulfur dichloride? A. SCl B. SCl2 C. S2Cl D. S2Cl2 E. S4Cl2
58. What is the correct formula for potassium dichromate? A. K2Cr2O7 B. K2(Cr2O7)2 C. K2CrO4 D. K2(CrO4)2 E. KCrO4
59. What is the formula for hypochlorous acid? A. HCl B. HClO C. HClO2 D. HClO3 E. HClO4
60. What is the correct name for H2SO4(aq)? A. sulfuric acid B. sulfide acid C. sulfurous acid D. hydrogen sulfate acid E. hydrogen sulfide acid
61. Using the laws of constant composition and the conservation of mass, complete the molecular picture of hydrogen molecules (circles) reacting with oxygen molecules (squares) to give water.
A.
B.
C.
D.
E.
62. Using the laws of constant composition and the conservation of mass, complete the molecular picture of hydrogen molecules (circles) reacting with chlorine molecules (squares) to give hydrogen chloride (HCl).
A.
B.
C.
D.
E.
63. BAC stands for: A. Breath Alcohol Concentration B. Blood Alcohol Concentration C. Brain Alcohol Concentration D. Blood Alcohol Consumption E. Bad Alcohol Correlation
64. Which of the following is a non-electrolyte in water? A. NaCl B. SF6 C. KNO3 D. MgS E. NH4Cl
Chapter 2--Atoms, Molecules, and Ions Key
1. Which of the following statements is/are postulates of Dalton’s atomic theory? 1. 2. 3.
Elements are composed of tiny particles called atoms. No atom is changed into another element in an ordinary chemical reaction. Compounds are formed when two or more atoms combine.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2 and 3 2. J.J. Thomson determined that electrons are small, negatively charged particles by A. bombarding gold foil with alpha particles. B. exposing photographic plates to radioactive uranium. C. deflecting cathode rays with electric and magnetic fields. D. converting cathode rays to electron particles using a fluorescent screen. E. decomposing neutrons into protons and electrons.
3. Which of the following statements is/are CORRECT? 1. 2. 3.
A neutron is an uncharged particle with a mass slightly greater than a proton. The nucleus of an atom has a negative charge. More than 99.9% of an atom's mass is concentrated in the electron cloud surrounding the nucleus.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2 and 3 4. Rank the subatomic particles from least to greatest mass. A. electrons = neutrons = protons B. electrons = protons < neutrons C. electrons < neutrons = protons D. electrons < protons < neutrons E. electrons < neutrons < protons
5. All of the following statements are true EXCEPT A. all atoms of a given element have the same mass number. B. for any neutral element, the number of electrons is equal to the number of protons. C. the mass number is the sum of the number of protons and neutrons. D. isotopes of atoms contain the same number of protons but a different number of neutrons. E. the atomic number equals the number of protons in an atom.
6. All atoms of the same element have the same number of ____. A. neutrons B. protons C. protons and neutrons D. electrons and neutrons E. protons, neutrons, and electrons
7. What is the mass number of an argon atom with 22 neutrons? A. 2 B. 18 C. 22 D. 40 E. 39.95
8. Which of the following atoms contains the fewest protons? A. 232Th B. 231Pa C. 245Pu D. 238U E. 232Pa
9. How many protons, neutrons, and electrons are in a silver atom with a mass number of 108? A. 47 protons, 47 neutrons, 61 electrons B. 47 protons, 61 neutrons, 47 electrons C. 61 protons, 47 neutrons, 47 electrons D. 47 protons, 108 neutrons, 47 electrons E. 61 protons, 108 neutrons, 61 electrons
10. F-20, a radioactive isotope of fluorine, has A. 9 protons, 10 neutrons, and 1 electron. B. 9 protons, 10 neutrons, and 9 electrons. C. 9 protons, 11 neutrons, and 9 electrons. D. 10 protons, 9 neutrons, and 1 electron. E. 10 protons, 10 neutrons, and 10 electrons.
11. Which two atoms below have the same number of neutrons? , A.
and
B.
and
C.
and
D.
and
E.
and
,
,
,
12. What is the atomic symbol for an element with 39 protons and 50 neutrons? A. B. C. D. E.
13. Which two of the ions below have the same number of electrons? ,
,
A.
and
B.
and
C.
and
D.
and
E.
and
,
,
14. Two isotopes of a given element will have the same number of ____, but a different number of ____ in their nucleus. A. protons, electrons B. electrons, protons C. protons, neutrons D. neutrons, protons E. electrons, neutrons
15. Which of the following are a pair of isotopes? A.
and
B.
and
C.
and
D.
and
E.
and
16. Which species has 63 neutrons? A. B. C. D. E. none of the above
17. Two isotopes of chlorine are found in nature, Cl-35 and Cl-37. The average mass of chlorine is 35.45 amu. The more abundant isotope of Cl has A. 17 protons, 17 electrons, and 18 neutrons. B. 17 protons, 17 electrons, and 18.45 neutrons. C. 17 protons, 17 electrons, and 20 neutrons. D. 18 protons, 18 electrons, and 17 neutrons. E. 19 protons, 19 electrons, and 16 neutrons.
18. What is the identity of A. zinc B. silver C. iridium D. cesium E. manganese
?
19. What is the symbol for an element which contains 57 neutrons and has a mass number of 101? A. Er B. Ru C. Md D. La E. Os
20. Rubidium has two naturally occurring isotopes. The average mass of Rb is 85.4678 amu. If 72.15% of Rb is found as Rb-85 (84.9117 amu), what is the mass of the other isotope? A. 0.56 amu B. 85.68 amu C. 86.91 amu D. 86.02 amu E. 83.47 amu
21. Silver has two stable isotopes with masses of 106.90509 amu and 108.9047 amu. The average molar mass of silver is 107.868 amu. What is the percent abundance of each isotope? A. 50.0% Ag-107 and 50.0% Ag-109 B. 51.8% Ag-107 and 48.2% Ag-109 C. 55.4% Ag-107 and 44.6% Ag-109 D. 48.2% Ag-107 and 51.8% Ag-109 E. 44.6% Ag-107 and 55.4% Ag-109
22. Gallium has an average atomic mass of 69.7 amu. In a typical sample, 60.4% of Ga exists as Ga-69 (68.9257 amu). What is the identity and the atomic mass of the other isotope? A.
; 70.9 amu
B.
; 70.9 amu
C.
; 70.9 amu
D.
; 71.9 amu
E.
; 71.9 amu
23. An element has three naturally occurring isotopes with the following abundances and masses:
abundance 78.99% 10.00% 11.01%
Determine the molar mass of the element.
A. 24.31 g/mol B. 24.98 g/mol C. 74.95 g/mol D. 2431 g/mol E. none of the above
mass (amu) 23.985042 24.985837 25.982593
24. The average molar mass of lithium is 6.941. A sample of lithium consists of two isotopes with masses of 6.01512 amu and 7.01600 amu. Determine the percent abundance of each isotope. A. 7.49% Li-6 and 92.51% Li-7 B. 8.45% Li-6 and 91.55% Li-7 C. 12.49% Li-6 and 87.51% Li-7 D. 91.55% Li-6 and 8.45% Li-7 E. 92.51% Li-6 and 7.49% Li-7
25. What is the mass (in grams) of a boron atom? A. 10.8 g B. 1.80 ´ 10-23 g C. 1.66 ´ 10-24 g D. 5.57 ´ 1022 g E. 1.54 ´ 10-25 g
26. Group 1 elements are also known as A. alkaline earth metals. B. alkali metals. C. chalcogens. D. halogens. E. noble gases.
27. How many nonmetallic elements are there in group 13? A. 0 B. 1 C. 2 D. 3 E. 4
28. Identify the halogen from period 4. A. Br B. I C. Kr D. Ar E. K
29. What element is in the fourth period in Group 3A? A. Sb B. Ga C. In D. Si E. Tl
30. Which group of three elements contains a nonmetal, a metal, and a metalloid? A. Li, Al, Si B. Na, Hg, I C. I, Hg, Si D. K, O, Br E. H, Al, N
31. Which three elements are likely to have similar chemical and physical properties? A. boron, silicon, and germanium B. sodium, magnesium, and aluminum C. sodium, potassium, and rubidium D. oxygen, sulfur, and chlorine E. carbon, nitrogen, and oxygen
32. Which group of three elements contains a transition metal, a halogen, and a noble gas? A. S, I, Cu B. Br, Kr, Ba C. Ar, Hg, Rn D. Ce, N, He E. Cu, I, Xe
33. How many elements are contained in period 2? A. 3 B. 8 C. 10 D. 18 E. 32
34. Of the naturally occurring elements in group 14, how many are nonmetals, metalloids, and metals? A. 0 nonmetals, 3 metalloids, and 2 metals B. 1 nonmetal, 2 metalloids, and 2 metals C. 2 nonmetals, 2 metalloids, and 1 metal D. 2 nonmetals, 1 metalloid, and 2 metals E. 3 nonmetals, 0 metalloids, and 2 metals
35. Which two of the following elements are abundant in the Earth's crust, but missing from the human body: O, Al, Si, Fe, C, N? A. O and Fe B. Si and C C. Al and Si D. O and N E. Fe and N
36. The formula of ethanol, CH3CH2OH, is an example of a(n) A. condensed formula. B. empirical formula. C. structural formula. D. ionic compound formula. E. mass spectrum.
37. Which element is most likely to form an ion with a -2 charge? A. K B. Mg C. P D. Br E. S
38. Which atom is most likely to form an ion with a +2 charge? A. scandium B. calcium C. aluminum D. oxygen E. fluorine
39. A strontium ion has ____ electrons. A. 35 B. 36 C. 37 D. 38 E. 39
40. For a nonmetal in Group 16 of the periodic table, the most common monatomic ion will have a charge of ____. A. –3 B. –2 C. –1 D. +1 E. +2
41. Identify the ions and their charges in Na2SO4. A. Na+, SO4B. Na+, SO42C. Na+, SO4D. Na2+, SO4E. Na2+, SO42-
42. Identify the ions and their charges in KH2PO4. A. K+, H+, P3-, O2B. K+, H2+, P3-, O8C. K+, H22+, P-1, O4-2 D. K+, H2PO4E. K+, H2+, PO43-
43. What is the correct formula for an ionic compound that contains magnesium ions and phosphide ions? A. MgP B. MgP2 C. Mg3P2 D. Mg3(PO4)2 E. Mg2P3
44. What is the correct formula for an ionic compound that contains aluminum ions and chloride ions? A. AlCl B. AlCl2 C. AlCl3 D. Al2Cl3 E. Al3Cl2
45. What are the values for x and y, respectively, in CaxHyPO4? A. 1 and 2 B. 2 and 1 C. 1 and 3 D. 2 and 2 E. 1 and 1
46. Sodium sulfate has the chemical formula Na2SO4. Based on this information, the formula for chromium(III) sulfate is ____. A. CrSO4 B. Cr(SO4)3 C. Cr2(SO4)3 D. Cr2SO4 E. Cr3(SO4)2
47. What is the correct name for MnS? A. manganese sulfide B. dimanganese sulfate C. dimanganese sulfide D. manganese(II) sulfate E. manganese(II) sulfide
48. What is the correct name for K3PO4? A. tripotassium phosphate B. potassium(I) monophosphorus tetraoxide C. potassium(I) phosphate D. potassium phosphate E. potassium phosphide
49. What is the correct name for TiCl4? A. monotitanium tetrachloride B. tetrachlorine titanate C. titanium tetrachlorine D. titanium(IV) tetrachloride E. titanium(IV) chloride
50. What is the correct name for Al2O3? A. alum B. aluminum trioxide C. aluminum ozinide D. aluminum oxide E. dialuminum trioxide
51. What is the correct formula for aluminum selenide? A. AlSe B. AlSe2 C. Al2Se D. Al2Se3 E. Al3Se2
52. What is the correct formula for iron(II) nitrate? A. Fe2(NO3)2 B. Fe2NO3 C. Fe(NO3)2 D. Fe3N2 E. FeNO3
53. What is the correct formula for barium perchlorate? A. BaClO4 B. BaClO3 C. Ba(ClO4)2 D. Ba(ClO3)2 E. Ba(ClO3)3
54. What is the correct name for N2O3? A. nitrogen oxide B. nitrogen(II) oxide C. nitrogen(III) oxide D. trioxygen dinitride E. dinitrogen trioxide
55. What is the correct name for PF5? A. phosphorus pentafluoride B. phosphorus(V) fluoride C. phosphorofluoride D. pentafluorophosphorus E. pentafluorophosphate
56. What is the correct name for CCl4? A. carbon chlorine B. tetracarbon chloride C. carbon tetrachloride D. carbon(IV) chloride E. tetrachlorocarbide
57. What is the correct formula for sulfur dichloride? A. SCl B. SCl2 C. S2Cl D. S2Cl2 E. S4Cl2
58. What is the correct formula for potassium dichromate? A. K2Cr2O7 B. K2(Cr2O7)2 C. K2CrO4 D. K2(CrO4)2 E. KCrO4
59. What is the formula for hypochlorous acid? A. HCl B. HClO C. HClO2 D. HClO3 E. HClO4
60. What is the correct name for H2SO4(aq)? A. sulfuric acid B. sulfide acid C. sulfurous acid D. hydrogen sulfate acid E. hydrogen sulfide acid
61. Using the laws of constant composition and the conservation of mass, complete the molecular picture of hydrogen molecules (circles) reacting with oxygen molecules (squares) to give water.
A.
B.
C.
D.
E.
62. Using the laws of constant composition and the conservation of mass, complete the molecular picture of hydrogen molecules (circles) reacting with chlorine molecules (squares) to give hydrogen chloride (HCl).
A.
B.
C.
D.
E.
63. BAC stands for: A. Breath Alcohol Concentration B. Blood Alcohol Concentration C. Brain Alcohol Concentration D. Blood Alcohol Consumption E. Bad Alcohol Correlation
64. Which of the following is a non-electrolyte in water? A. NaCl B. SF6 C. KNO3 D. MgS E. NH4Cl
Chapter 3--Mass Relations in Chemistry; Stoichiometry 1. If a penny has a mass of 2.51 g, what is the mass of 1.00 millimoles of pennies? (1 millimole = 1 ´ 10-3 mole) A. 1.51 ´ 1021 g B. 2.51 ´ 103 g C. 6.02 ´ 1020 g D. 2.51 ´ 10-3 g E. 2.38 ´ 1022 g
2. What is the mass of 0.71 mol Na? A. 1.2 ´ 10–24 g B. 12 g C. 16 g D. 0.031 g E. 32 g
3. The molar mass of boron is 10.81 g/mole. What is the mass of a single boron atom? A. 1.661 ´ 10-24 g B. 1.795 ´ 10-23 g C. 6.510 ´ 1024 g D. 1.536 ´ 10-25 g E. 1.081 ´ 101 g
4. The molar mass of nitrogen (N2) is 28.0 g/mole. What is the mass of a single nitrogen atom? A. 2.32 ´ 10-23 g B. 4.65 ´ 10-23 g C. 9.30 ´ 10-23 g D. 4.30 ´ 10-22 g E. 8.43 ´ 10-22 g
5. The mass of a single molecule of elemental phosphorus is 2.057 ´ 10-22 g. How many atoms combine to form a single molecule of elemental phosphorus? A. 1 B. 2 C. 4 D. 6 E. 10
6. The molar mass of platinum is 195.08 g/mol. What is the mass of 1.00 ´ 102 Pt atoms? A. 8.51 ´ 10–25 g B. 3.24 ´ 10–24 g C. 1.67 ´ 10–22 g D. 3.24 ´ 10–22 g E. 3.24 ´ 10–20 g
7. The mass of a single atom of chlorine atom is 5.887 ´ 10-23 grams. Which is a correct method for determining the molar mass of elemental chlorine, Cl2?
A.
B.
C.
D.
E.
8. A 1.45 g sample of chromium contains ____ atoms. A. 1.25 ´ 1022 B. 1.68 ´ 1022 C. 8.73 ´ 1023 D. 2.16 ´ 1025 E. 4.54 ´ 1025
9. You have a 5.0 g sample of each of the following elements: Ra, Rb, Rh, Rn, Ru. Which sample contains the most atoms? A. Ra B. Rb C. Rh D. Rn E. Ru
10. Which of the following samples contains the largest number of atoms? A. 2.0 moles of H3PO4 B. 3.0 moles of H2SO3 C. 4.0 moles of HNO3 D. 6.0 moles of HClO E. 8.0 moles of HBr
11. Which of the following samples contains the largest number of hydrogen atoms? A. 2.0 moles of C6H16 B. 3.0 moles of C3H8 C. 4.0 moles of C3H6 D. 6.0 moles of C2H4 E. 8.0 moles of C2H2
12. What mass of chlorine is present in 5.00 grams of carbon tetrachloride (CCl4)? A. 0.130 g B. 1.15 g C. 0.564 g D. 4.61 g E. 0.922 g
13. What is the mass in grams of 0.338 mol of glucose (C6H12O6)? A. 0.00188 g B. 0.0164 g C. 1.88 g D. 53.3 g E. 60.9 g
14. What is the mass in grams of 0.362 moles barium chloride (BaCl2)? A. 0.362 g B. 0.00174 g C. 75.4 g D. 208 g E. 575 g
15. Which is a correct method for determining the mass in grams of 0.190 mol silver?
A.
B.
C.
D. E. none of the above
16. How many moles of HCl are present in 0.098 grams of HCl? A. 0.00083 mol B. 0.0027 mol C. 0.28 mol D. 3.6 mol E. 380 mol
17. Which is a correct method for calculating the moles of calcium carbonate present in 132 grams of calcium carbonate?
A.
B.
C.
D. E. none of the above
18. How many atoms are present in 5.00 grams of iron? A. 8.95 ´ 10-2 atoms B. 3.36 ´ 1018 atoms C. 1.88 ´ 1021 atoms D. 5.39 ´ 1022 atoms E. 3.36 ´ 1026 atoms
19. How many hydrogen atoms are present in 1.0 g of NH3? A. 0.059 atoms B. 0.18 atoms C. 3.5 ´ 1022 atoms D. 1.1 ´ 1023 atoms E. 1.2 ´ 1022 atoms
20. Which is a correct method for determining the total number of atoms in 123 grams of sulfur trioxide (SO3)?
A.
B.
C.
D.
E.
21. What mass of oxygen is present in 10.0 grams of potassium nitrate (KNO3)? A. 4.75 g B. 5.43 g C. 6.39 g D. 8.00 g E. 9.17 g
22. The molarity of a solution is defined as A. the moles of solute per liter of solution. B. the moles of solute per kilogram of solution. C. the moles of solute per kilogram of solvent. D. the mass of solute (in grams) per liter of solution. E. the mass of solute (in grams) per liter of solvent.
23. What is the maximum volume of 0.25 M KCl(aq) that can be prepared from 75 g KCl(s)? A. 0.33 L B. 1.0 L C. 3.0 L D. 4.0 L E. 19 L
24. If 8.19 g KIO3 is dissolved in enough water to make 500.0 mL of solution, what is the molarity of the potassium iodate solution? The molar mass of KIO3 is 214 g/mol. A. 1.64 ´ 10–2 M B. 1.91 ´ 10–2 M C. 7.65 ´ 10–2 M D. 3.51 M E. 16.4 M
25. How many liters of 0.2805 M C6H12O6(aq) contain 1.000 g of C6H12O6? A. 0.001557 L B. 0.01979 L C. 0.2805 L D. 3.565 L E. 50.5 L
26. If 5.15 g Fe(NO3)3 is dissolved in enough water to make exactly 150.0 mL of solution, what is the molar concentration of nitrate ion? A. 0.00319 M B. 0.0343 M C. 0.142 M D. 0.313 M E. 0.426 M
27. A mass of 12.0 g of calcium chloride is diluted to a volume of 250 mL in a volumetric flask. Which of the equations below is a correct method for determining the chloride ion concentration?
A.
B.
C.
D.
E.
28. What is the mass of sodium iodide in 50.0 mL of 2.63 ´ 10–2 M NaI(aq)? A. 0.00132 g B. 0.00877 g C. 0.0788 g D. 0.197 g E. 78.8 g
29. If 25.00 mL of 4.50 M NaOH(aq) is diluted with water to a volume of 750.0 mL, what is the molarity of the diluted NaOH(aq)? A. 0.0333 M B. 0.150 M C. 0.155 M D. 6.67 M E. 1.35 ´ 103 M
30. What volume of 0.15 M HCl(aq) must be diluted to make 2.0 L of 0.050 M HCl(aq)? A. 0.015 L B. 0.10 L C. 0.30 L D. 0.67 L E. 6.0 L
31. What is the mass percent of each element in sulfuric acid, H2SO4? A. 2.055% H, 32.69% S, 65.25% O B. 1.028% H, 32.69% S, 66.28% O C. 28.57% H, 14.29% S, 57.17% O D. 1.028% H, 33.72% S, 65.25% O E. 2.016% H, 32.07% S, 65.91% O
32. What is the percent composition of iron(II) sulfate hexahydrate? A. 4.2% Fe; 4.2% S; 41.6% O; 50.0% H B. 16.7% Fe; 16.7%S; 66.6% O C. 21.5% Fe; 12.3%S; 24.6% O; 41.6% H D. 21.5% Fe; 12.3%S; 61.5% O; 4.7% H E. 36.8% Fe; 21.1%S; 42.1% O
33. Which is a correct method for determining the mass of carbon present in 0.132 grams of propane (C3H8)?
A.
B.
C.
D.
E.
34. Nitrogen and oxygen form an extensive series of oxides with the general formula NxOy. What is the empirical formula for an oxide that contains 30.44% by mass nitrogen? A. N2O B. NO C. NO2 D. N2O3 E. N2O5
35. Beryl is a mineral which contains 5.03% Be, 10.04% Al, 31.35% Si, and 53.58% O. What is the simplest formula for beryl? A. BeAl(SiO3)2 B. BeAl(SiO3)3 C. Be3(AlSiO3)2 D. Be3Al2(SiO3)6 E. Be4Al(SiO3)8
36. A molecule is found to contain 47.35% by mass C, 10.60% by mass H, and 42.05% by mass O. What is the empirical formula for this molecule? A. C2H6O B. C3H4O C. C3H8O2 D. C4H6O2 E. C4H8O3
37. Polyethylene is a polymer consisting of only carbon and hydrogen. If 2.300 g of the polymer is burned in oxygen it produces 2.955 g H2O and 7.217 g CO2. What is the empirical formula of polyethylene? A. CH B. CH2 C. C2H3 D. C5H8 E. C7H8
38. Isopentyl acetate, a molecule composed of C, H, and O, smells like bananas. Combustion analysis of 1.750 grams of this molecule yields 1.695 g H2O and 4.142 g CO2. What is the simplest formula for isopentyl acetate? A. C7H14O2 B. C7H7O4 C. C8H10O3 D. C8H16O E. C9H6O
39. Combustion analysis of 0.800 grams of an unknown hydrocarbon yields 2.613 g CO2 and 0.778 g H2O. What is the percent composition of the hydrocarbon? A. 66.6% C; 33.4% H B. 82.3% C; 17.7% H C. 89.1% C; 10.9% H D. 92.4% C; 7.60% H E. not enough information given to solve the problem
40. Soft drink bottles are made of polyethylene terephthalate (PET), a polymer composed of carbon, hydrogen, and oxygen. If 2.8880 g PET is burned in oxygen it produces 1.0000 g H2O and 6.1058 g CO2. What is the empirical formula of PET? A. CHO B. CH7O5 C. C5H7O D. C8H10O E. C10H8O5
41. An oxide of nitrogen contains 63.1% oxygen and has a molar mass of 76.0 g/mol. What is the molecular formula for this compound? A. N2O B. NO C. NO2 D. N2O3 E. N2O5
42. Nitroglycerin decomposes violently according to the balanced chemical equation below. 2C3H5(NO3)3(l) ® 3N2(g) + 1/2O2(g) + 6CO2(g) + 5H2O(g) Which of the following statements concerning this reaction is/are CORRECT? 1. 2. 3.
Two moles of nitroglycerine will produce three moles of nitrogen and five moles of water. Four molecules of nitroglycerine will produced one molecule of oxygen and twelve molecules of carbon dioxide. Six grams of nitroglycerine will produce nine grams of nitrogen and fifteen grams of water.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 43. Hydrazine, a fuel used to power rocket engines, is a product of the reaction between ammonia and bleach. Balance the equation for the reaction. aNH3(aq) + bOCl-(aq) ® cN2H4(l) + dCl-(aq) + eH2O(l) A. a = 2, b = 1, c = 1, d = 1, e = 1 B. a = 2, b = 2, c = 1, d = 2, e = 2 C. a = 2, b = 1, c = 2, d = 1, e = 2 D. a = 4, b = 3, c = 1, d = 3, e = 1 E. a = 4, b = 3, c = 2, d = 3, e = 2
44. What is the balanced chemical equation for the complete combustion of benzoic acid, C6H5CO2H, to form carbon dioxide and water? A. C6H5CO2H(s) ® 6C(s) + CO2(g) + 3H2(g) B. C6H5CO2H(s) ® 7CO2(g) + 3H2O(g) C. C6H5CO2H(s) + O2(g) ® CO2(g) + H2O(g) D. C6H5CO2H(s) + 8O2(g) ® 7CO2(g) + 3H2O(g) E. 2C6H5CO2H(s) + 15O2(g) ® 14CO2(g) + 6H2O(g)
45. Aluminum reacts in air with oxygen to form aluminum oxide. Which of the reactions below is correct and properly balanced? A. Al(s) + O(g) ® AlO(s) B. 2Al(s) + O2(g) ® 2AlO(s) C. 3Al(s) + O2(g) ® Al3O2(s) D. 4Al(s) + O2(g) ® 2Al2O(s) E. 4Al(s) + 3O2(g) ® 2Al2O3(s)
46. If 16.4 g of oxygen gas react with excess hydrogen, what mass of water is produced? 2H2(g) + O2(g) ® 2H2O(g) A. 9.23 g B. 18.5 g C. 20.4 g D. 23.9 g E. 36.9 g
47. Iron reacts with hydrochloric acid to produce iron(II) chloride and hydrogen gas. Fe(s) + 2 HCl(aq) ® FeCl2(aq) + H2(g) What mass of H2(g) is produced from the reaction of 5.2 g Fe(s) with excess hydrochloric acid? A. 0.094 g B. 0.19 g C. 5.2 g D. 6.8 g E. 1.4 ´ 102 g
48. How many moles of ammonia can be made by reacting 7.0 mol of N2 with 4.0 mol of H2? N2(g) + 3H2(g) ® 2NH3(g) A. 2.7 mol B. 4.0 mol C. 7.0 mol D. 11 mol E. 14 mol
49. The reaction of coal and water at a high temperature produces a mixture of hydrogen and carbon monoxide gases. This mixture is known as synthesis gas (or syngas). What mass of carbon monoxide can be formed from the reaction of 71.3 g of carbon with excess water? C(s) + H2O(g) ® H2(g) + CO(g) A. 5.94 g B. 12.0 g C. 71.3 g D. 107 g E. 166 g
50. What mass of oxygen will react with 2.64 g of magnesium? 2Mg(s) + O2(g) ® MgO(s) A. 0.487 g B. 1.00 g C. 1.26 g D. 1.74 g E. 3.47 g
51. The compound P4S3 is used in matches. It reacts with oxygen to produce P4O10 and SO2. The unbalanced chemical equation is shown below. P4S3(s) + O2(g) ® P4O10(s) + SO2(g) What mass of SO2 is produced from the combustion of 0.331 g P4S3? A. 0.00150 g B. 0.00451 g C. 0.0321g D. 0.0964 g E. 0.289 g
52. What mass of carbon dioxide can be made by reacting 1.56 grams of sodium bicarbonate with 0.687 grams of hydrochloric acid? NaHCO3(s) + H+(aq) ® CO2(g) + H2O(l) + Na+(aq) A. 2.25 g B. 2.98 g C. 0.817 g D. 0.829 g E. 11.4 g
53. Nitric oxide is made from the oxidation of ammonia. What mass of nitric oxide can be made from the reaction of 8.00 g NH3 with 17.0 g O2? 4NH3(g) + 5O2(g) ® 4NO(g) + 6H2O(g) A. 4.54 g B. 12.8 g C. 14.1 g D. 15.9 g E. 25.0 g
54. A mass of 8.15 g C2H4(g) reacts with excess oxygen. If 16.2 g CO2(g) is collected, what is the percent yield of the reaction? C2H4(g) + 3O2(g) ® 2CO2(g) + 2H2O(g) A. 25.6% B. 31.7% C. 41.0% D. 57.1% E. 63.3%
55. A mass of 4.00 g of H2(g) reacts with 2.00 g of O2(g). If 1.94 g of H2O(l) is collected, what is the percent yield of the reaction? 2H2(g) + O2(g) ® 2H2O(l) A. 5.4 % B. 49 % C. 32 % D. 86 % E. 97 %
56. Under certain conditions the reaction of ammonia with excess oxygen will produce a 24.8% yield of NO. What mass of NH3 must react with excess oxygen to yield 12.5 g NO? 4NH3(g) + 5O2(g) ® 4NO(g) + 6H2O(g) A. 1.76 g B. 7.10 g C. 28.6 g D. 50.4 g E. 88.8 g
57. A 3.592 g sample of hydrated magnesium bromide, MgBr2xH2O, is dried in an oven. When the anhydrous salt is removed from the oven, its mass is 2.263 g. What is the value of x? A. 1 B. 3 C. 6 D. 8 E. 12
58. Which of the balanced chemical equations is consistent with the following pictoral representation of a chemical reaction?
A. 2SO2 + O2 ® 2SO3 B. H2 + I2 ® 2HI C. 2H2 + O2 ® 2H2O D. 2N2 + 3H2 ® 2NH3 E. none of the above
59. What hydrate is sometimes referred to as "the ice that burns"? A. copper(II) sulfate pentahydrate B. carbon dioxide hydrate C. methane hydrate D. cobalt(II) chloride pentahydrate E. none of the above
60. Chlorophyll, the substance responsible for the green color of leaves, has one magnesium atom per chlorophyll molecule and contains 2.72% magnesium by mass. What is the molar mass of chlorophyll? A. 24.3 g/mol B. 20.2 g/mol C. 2020 g/mol D. 8.94 g/mol E. 894 g/mol
Chapter 3--Mass Relations in Chemistry; Stoichiometry Key 1. If a penny has a mass of 2.51 g, what is the mass of 1.00 millimoles of pennies? (1 millimole = 1 ´ 10-3 mole) A. 1.51 ´ 1021 g B. 2.51 ´ 103 g C. 6.02 ´ 1020 g D. 2.51 ´ 10-3 g E. 2.38 ´ 1022 g
2. What is the mass of 0.71 mol Na? A. 1.2 ´ 10–24 g B. 12 g C. 16 g D. 0.031 g E. 32 g
3. The molar mass of boron is 10.81 g/mole. What is the mass of a single boron atom? A. 1.661 ´ 10-24 g B. 1.795 ´ 10-23 g C. 6.510 ´ 1024 g D. 1.536 ´ 10-25 g E. 1.081 ´ 101 g
4. The molar mass of nitrogen (N2) is 28.0 g/mole. What is the mass of a single nitrogen atom? A. 2.32 ´ 10-23 g B. 4.65 ´ 10-23 g C. 9.30 ´ 10-23 g D. 4.30 ´ 10-22 g E. 8.43 ´ 10-22 g
5. The mass of a single molecule of elemental phosphorus is 2.057 ´ 10-22 g. How many atoms combine to form a single molecule of elemental phosphorus? A. 1 B. 2 C. 4 D. 6 E. 10
6. The molar mass of platinum is 195.08 g/mol. What is the mass of 1.00 ´ 102 Pt atoms? A. 8.51 ´ 10–25 g B. 3.24 ´ 10–24 g C. 1.67 ´ 10–22 g D. 3.24 ´ 10–22 g E. 3.24 ´ 10–20 g
7. The mass of a single atom of chlorine atom is 5.887 ´ 10-23 grams. Which is a correct method for determining the molar mass of elemental chlorine, Cl2?
A.
B.
C.
D.
E.
8. A 1.45 g sample of chromium contains ____ atoms. A. 1.25 ´ 1022 B. 1.68 ´ 1022 C. 8.73 ´ 1023 D. 2.16 ´ 1025 E. 4.54 ´ 1025
9. You have a 5.0 g sample of each of the following elements: Ra, Rb, Rh, Rn, Ru. Which sample contains the most atoms? A. Ra B. Rb C. Rh D. Rn E. Ru
10. Which of the following samples contains the largest number of atoms? A. 2.0 moles of H3PO4 B. 3.0 moles of H2SO3 C. 4.0 moles of HNO3 D. 6.0 moles of HClO E. 8.0 moles of HBr
11. Which of the following samples contains the largest number of hydrogen atoms? A. 2.0 moles of C6H16 B. 3.0 moles of C3H8 C. 4.0 moles of C3H6 D. 6.0 moles of C2H4 E. 8.0 moles of C2H2
12. What mass of chlorine is present in 5.00 grams of carbon tetrachloride (CCl4)? A. 0.130 g B. 1.15 g C. 0.564 g D. 4.61 g E. 0.922 g
13. What is the mass in grams of 0.338 mol of glucose (C6H12O6)? A. 0.00188 g B. 0.0164 g C. 1.88 g D. 53.3 g E. 60.9 g
14. What is the mass in grams of 0.362 moles barium chloride (BaCl2)? A. 0.362 g B. 0.00174 g C. 75.4 g D. 208 g E. 575 g
15. Which is a correct method for determining the mass in grams of 0.190 mol silver?
A.
B.
C.
D. E. none of the above
16. How many moles of HCl are present in 0.098 grams of HCl? A. 0.00083 mol B. 0.0027 mol C. 0.28 mol D. 3.6 mol E. 380 mol
17. Which is a correct method for calculating the moles of calcium carbonate present in 132 grams of calcium carbonate?
A.
B.
C.
D. E. none of the above
18. How many atoms are present in 5.00 grams of iron? A. 8.95 ´ 10-2 atoms B. 3.36 ´ 1018 atoms C. 1.88 ´ 1021 atoms D. 5.39 ´ 1022 atoms E. 3.36 ´ 1026 atoms
19. How many hydrogen atoms are present in 1.0 g of NH3? A. 0.059 atoms B. 0.18 atoms C. 3.5 ´ 1022 atoms D. 1.1 ´ 1023 atoms E. 1.2 ´ 1022 atoms
20. Which is a correct method for determining the total number of atoms in 123 grams of sulfur trioxide (SO3)?
A.
B.
C.
D.
E.
21. What mass of oxygen is present in 10.0 grams of potassium nitrate (KNO3)? A. 4.75 g B. 5.43 g C. 6.39 g D. 8.00 g E. 9.17 g
22. The molarity of a solution is defined as A. the moles of solute per liter of solution. B. the moles of solute per kilogram of solution. C. the moles of solute per kilogram of solvent. D. the mass of solute (in grams) per liter of solution. E. the mass of solute (in grams) per liter of solvent.
23. What is the maximum volume of 0.25 M KCl(aq) that can be prepared from 75 g KCl(s)? A. 0.33 L B. 1.0 L C. 3.0 L D. 4.0 L E. 19 L
24. If 8.19 g KIO3 is dissolved in enough water to make 500.0 mL of solution, what is the molarity of the potassium iodate solution? The molar mass of KIO3 is 214 g/mol. A. 1.64 ´ 10–2 M B. 1.91 ´ 10–2 M C. 7.65 ´ 10–2 M D. 3.51 M E. 16.4 M
25. How many liters of 0.2805 M C6H12O6(aq) contain 1.000 g of C6H12O6? A. 0.001557 L B. 0.01979 L C. 0.2805 L D. 3.565 L E. 50.5 L
26. If 5.15 g Fe(NO3)3 is dissolved in enough water to make exactly 150.0 mL of solution, what is the molar concentration of nitrate ion? A. 0.00319 M B. 0.0343 M C. 0.142 M D. 0.313 M E. 0.426 M
27. A mass of 12.0 g of calcium chloride is diluted to a volume of 250 mL in a volumetric flask. Which of the equations below is a correct method for determining the chloride ion concentration?
A.
B.
C.
D.
E.
28. What is the mass of sodium iodide in 50.0 mL of 2.63 ´ 10–2 M NaI(aq)? A. 0.00132 g B. 0.00877 g C. 0.0788 g D. 0.197 g E. 78.8 g
29. If 25.00 mL of 4.50 M NaOH(aq) is diluted with water to a volume of 750.0 mL, what is the molarity of the diluted NaOH(aq)? A. 0.0333 M B. 0.150 M C. 0.155 M D. 6.67 M E. 1.35 ´ 103 M
30. What volume of 0.15 M HCl(aq) must be diluted to make 2.0 L of 0.050 M HCl(aq)? A. 0.015 L B. 0.10 L C. 0.30 L D. 0.67 L E. 6.0 L
31. What is the mass percent of each element in sulfuric acid, H2SO4? A. 2.055% H, 32.69% S, 65.25% O B. 1.028% H, 32.69% S, 66.28% O C. 28.57% H, 14.29% S, 57.17% O D. 1.028% H, 33.72% S, 65.25% O E. 2.016% H, 32.07% S, 65.91% O
32. What is the percent composition of iron(II) sulfate hexahydrate? A. 4.2% Fe; 4.2% S; 41.6% O; 50.0% H B. 16.7% Fe; 16.7%S; 66.6% O C. 21.5% Fe; 12.3%S; 24.6% O; 41.6% H D. 21.5% Fe; 12.3%S; 61.5% O; 4.7% H E. 36.8% Fe; 21.1%S; 42.1% O
33. Which is a correct method for determining the mass of carbon present in 0.132 grams of propane (C3H8)?
A.
B.
C.
D.
E.
34. Nitrogen and oxygen form an extensive series of oxides with the general formula NxOy. What is the empirical formula for an oxide that contains 30.44% by mass nitrogen? A. N2O B. NO C. NO2 D. N2O3 E. N2O5
35. Beryl is a mineral which contains 5.03% Be, 10.04% Al, 31.35% Si, and 53.58% O. What is the simplest formula for beryl? A. BeAl(SiO3)2 B. BeAl(SiO3)3 C. Be3(AlSiO3)2 D. Be3Al2(SiO3)6 E. Be4Al(SiO3)8
36. A molecule is found to contain 47.35% by mass C, 10.60% by mass H, and 42.05% by mass O. What is the empirical formula for this molecule? A. C2H6O B. C3H4O C. C3H8O2 D. C4H6O2 E. C4H8O3
37. Polyethylene is a polymer consisting of only carbon and hydrogen. If 2.300 g of the polymer is burned in oxygen it produces 2.955 g H2O and 7.217 g CO2. What is the empirical formula of polyethylene? A. CH B. CH2 C. C2H3 D. C5H8 E. C7H8
38. Isopentyl acetate, a molecule composed of C, H, and O, smells like bananas. Combustion analysis of 1.750 grams of this molecule yields 1.695 g H2O and 4.142 g CO2. What is the simplest formula for isopentyl acetate? A. C7H14O2 B. C7H7O4 C. C8H10O3 D. C8H16O E. C9H6O
39. Combustion analysis of 0.800 grams of an unknown hydrocarbon yields 2.613 g CO2 and 0.778 g H2O. What is the percent composition of the hydrocarbon? A. 66.6% C; 33.4% H B. 82.3% C; 17.7% H C. 89.1% C; 10.9% H D. 92.4% C; 7.60% H E. not enough information given to solve the problem
40. Soft drink bottles are made of polyethylene terephthalate (PET), a polymer composed of carbon, hydrogen, and oxygen. If 2.8880 g PET is burned in oxygen it produces 1.0000 g H2O and 6.1058 g CO2. What is the empirical formula of PET? A. CHO B. CH7O5 C. C5H7O D. C8H10O E. C10H8O5
41. An oxide of nitrogen contains 63.1% oxygen and has a molar mass of 76.0 g/mol. What is the molecular formula for this compound? A. N2O B. NO C. NO2 D. N2O3 E. N2O5
42. Nitroglycerin decomposes violently according to the balanced chemical equation below. 2C3H5(NO3)3(l) ® 3N2(g) + 1/2O2(g) + 6CO2(g) + 5H2O(g) Which of the following statements concerning this reaction is/are CORRECT? 1. 2. 3.
Two moles of nitroglycerine will produce three moles of nitrogen and five moles of water. Four molecules of nitroglycerine will produced one molecule of oxygen and twelve molecules of carbon dioxide. Six grams of nitroglycerine will produce nine grams of nitrogen and fifteen grams of water.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 43. Hydrazine, a fuel used to power rocket engines, is a product of the reaction between ammonia and bleach. Balance the equation for the reaction. aNH3(aq) + bOCl-(aq) ® cN2H4(l) + dCl-(aq) + eH2O(l) A. a = 2, b = 1, c = 1, d = 1, e = 1 B. a = 2, b = 2, c = 1, d = 2, e = 2 C. a = 2, b = 1, c = 2, d = 1, e = 2 D. a = 4, b = 3, c = 1, d = 3, e = 1 E. a = 4, b = 3, c = 2, d = 3, e = 2
44. What is the balanced chemical equation for the complete combustion of benzoic acid, C6H5CO2H, to form carbon dioxide and water? A. C6H5CO2H(s) ® 6C(s) + CO2(g) + 3H2(g) B. C6H5CO2H(s) ® 7CO2(g) + 3H2O(g) C. C6H5CO2H(s) + O2(g) ® CO2(g) + H2O(g) D. C6H5CO2H(s) + 8O2(g) ® 7CO2(g) + 3H2O(g) E. 2C6H5CO2H(s) + 15O2(g) ® 14CO2(g) + 6H2O(g)
45. Aluminum reacts in air with oxygen to form aluminum oxide. Which of the reactions below is correct and properly balanced? A. Al(s) + O(g) ® AlO(s) B. 2Al(s) + O2(g) ® 2AlO(s) C. 3Al(s) + O2(g) ® Al3O2(s) D. 4Al(s) + O2(g) ® 2Al2O(s) E. 4Al(s) + 3O2(g) ® 2Al2O3(s)
46. If 16.4 g of oxygen gas react with excess hydrogen, what mass of water is produced? 2H2(g) + O2(g) ® 2H2O(g) A. 9.23 g B. 18.5 g C. 20.4 g D. 23.9 g E. 36.9 g
47. Iron reacts with hydrochloric acid to produce iron(II) chloride and hydrogen gas. Fe(s) + 2 HCl(aq) ® FeCl2(aq) + H2(g) What mass of H2(g) is produced from the reaction of 5.2 g Fe(s) with excess hydrochloric acid? A. 0.094 g B. 0.19 g C. 5.2 g D. 6.8 g E. 1.4 ´ 102 g
48. How many moles of ammonia can be made by reacting 7.0 mol of N2 with 4.0 mol of H2? N2(g) + 3H2(g) ® 2NH3(g) A. 2.7 mol B. 4.0 mol C. 7.0 mol D. 11 mol E. 14 mol
49. The reaction of coal and water at a high temperature produces a mixture of hydrogen and carbon monoxide gases. This mixture is known as synthesis gas (or syngas). What mass of carbon monoxide can be formed from the reaction of 71.3 g of carbon with excess water? C(s) + H2O(g) ® H2(g) + CO(g) A. 5.94 g B. 12.0 g C. 71.3 g D. 107 g E. 166 g
50. What mass of oxygen will react with 2.64 g of magnesium? 2Mg(s) + O2(g) ® MgO(s) A. 0.487 g B. 1.00 g C. 1.26 g D. 1.74 g E. 3.47 g
51. The compound P4S3 is used in matches. It reacts with oxygen to produce P4O10 and SO2. The unbalanced chemical equation is shown below. P4S3(s) + O2(g) ® P4O10(s) + SO2(g) What mass of SO2 is produced from the combustion of 0.331 g P4S3? A. 0.00150 g B. 0.00451 g C. 0.0321g D. 0.0964 g E. 0.289 g
52. What mass of carbon dioxide can be made by reacting 1.56 grams of sodium bicarbonate with 0.687 grams of hydrochloric acid? NaHCO3(s) + H+(aq) ® CO2(g) + H2O(l) + Na+(aq) A. 2.25 g B. 2.98 g C. 0.817 g D. 0.829 g E. 11.4 g
53. Nitric oxide is made from the oxidation of ammonia. What mass of nitric oxide can be made from the reaction of 8.00 g NH3 with 17.0 g O2? 4NH3(g) + 5O2(g) ® 4NO(g) + 6H2O(g) A. 4.54 g B. 12.8 g C. 14.1 g D. 15.9 g E. 25.0 g
54. A mass of 8.15 g C2H4(g) reacts with excess oxygen. If 16.2 g CO2(g) is collected, what is the percent yield of the reaction? C2H4(g) + 3O2(g) ® 2CO2(g) + 2H2O(g) A. 25.6% B. 31.7% C. 41.0% D. 57.1% E. 63.3%
55. A mass of 4.00 g of H2(g) reacts with 2.00 g of O2(g). If 1.94 g of H2O(l) is collected, what is the percent yield of the reaction? 2H2(g) + O2(g) ® 2H2O(l) A. 5.4 % B. 49 % C. 32 % D. 86 % E. 97 %
56. Under certain conditions the reaction of ammonia with excess oxygen will produce a 24.8% yield of NO. What mass of NH3 must react with excess oxygen to yield 12.5 g NO? 4NH3(g) + 5O2(g) ® 4NO(g) + 6H2O(g) A. 1.76 g B. 7.10 g C. 28.6 g D. 50.4 g E. 88.8 g
57. A 3.592 g sample of hydrated magnesium bromide, MgBr2xH2O, is dried in an oven. When the anhydrous salt is removed from the oven, its mass is 2.263 g. What is the value of x? A. 1 B. 3 C. 6 D. 8 E. 12
58. Which of the balanced chemical equations is consistent with the following pictoral representation of a chemical reaction?
A. 2SO2 + O2 ® 2SO3 B. H2 + I2 ® 2HI C. 2H2 + O2 ® 2H2O D. 2N2 + 3H2 ® 2NH3 E. none of the above
59. What hydrate is sometimes referred to as "the ice that burns"? A. copper(II) sulfate pentahydrate B. carbon dioxide hydrate C. methane hydrate D. cobalt(II) chloride pentahydrate E. none of the above
60. Chlorophyll, the substance responsible for the green color of leaves, has one magnesium atom per chlorophyll molecule and contains 2.72% magnesium by mass. What is the molar mass of chlorophyll? A. 24.3 g/mol B. 20.2 g/mol C. 2020 g/mol D. 8.94 g/mol E. 894 g/mol
Chapter 4--Reactions in Aqueous Solution 1. Which of the following statements concerning aqueous NaCl is/are CORRECT? 1. 2. 3.
Sodium chloride is a strong electrolyte. Sodium chloride is insoluble in water. NaCl(aq) consists of neutral NaCl molecules dissolved in water.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 2. Which of the following pairs of aqueous solutions will produce a precipitate when mixed? A. NaCl(aq) and Ca(NO3)2(aq) B. KCl(aq) and NaOH(aq) C. Cu(NO3)2(aq) and AgNO3(aq) D. ZnSO4(aq) and LiCl(aq) E. K2SO4(aq) and Pb(ClO4)2(aq)
3. Which of the following pairs of aqueous solutions will NOT produce a precipitate when mixed? A. AgNO3(aq) and KI(aq) B. Cu(NO3)2(aq) and MgCl2(aq) C. AlCl3(aq) and KOH(aq) D. BaBr2(aq) and Na2SO4(aq) E. AgNO3(aq) and NaCl(aq)
4. Precipitation reactions occur A. when group 1 cations are mixed with group 17 anions. B. when insoluble reactants are mixed. C. when ionic compounds react to form non-ionic products. D. predominantly with halide salts. E. when soluble ionic reactants combine to form insoluble products.
5. Which of the following statements is/are correct? 1. 2. 3.
Most ionic compounds containing nitrate ion are soluble in water. Most ionic compounds containing sulfate ion are insoluble in water. Most ionic compounds containing carbonate ion are soluble in water.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 6. Which of the following statements is/are correct? 1. 2. 3.
Most ionic compounds containing phosphate ion are insoluble in water. Most ionic compounds containing potassium ion are insoluble in water. Most ionic compounds containing hydroxide ion are soluble in water.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 7. If aqueous solutions of sodium sulfate and barium chloride are mixed, a white precipitate forms. What is the identity of the precipitate? A. Na2Ba B. NaCl2 C. NaCl D. BaSO4 E. Ba2SO4
8. If aqueous solutions of nickel(II) chloride and potassium phosphate are mixed, which insoluble compound will form? A. Ni3(PO4)2 B. NiCl2 C. KCl D. K2Ni E. K3PO4
9. Which of the following ionic compounds is not soluble in water? A. K2CO3 B. AgNO3 C. CaBr2 D. CsI E. PbI2
10. Which of the following ionic compounds are likely to be soluble in water: Mg(OH)2, Pb(NO3)2, AgI, Na2CO3, and Cu3(PO4)3? A. Na2CO3 only B. Pb(NO3)2 and Na2CO3 C. Mg(OH)2, Na2CO3, and Cu3(PO4)3 D. Mg(OH)2, AgI, and Cu3(PO4)3 E. Pb(NO3)2, AgI, Na2CO3, and Cu3(PO4)3
11. Which of the following compounds is least soluble in water? A. Fe(OH)3. B. Sr(OH)2. C. Al(NO3)3. D. KI. E. CuCl2.
12. If an aqueous solution of ____ is added to a mixture of Pb2+ and Ba2+, the lead ion will precipitate, but the barium ion will remain in solution. A. NaOH B. Na2SO4 C. K3PO4 D. KCO3 E. Ca(CH3CO2)2
13. What reaction occurs when aqueous solutions of sodium hydroxide and copper(II) bromide are mixed? A. Na+(aq) + OH-(aq) + Cu2+(aq) + 2Br-(aq) ® NaCu(s) + Br2OH(aq) B. Na+(aq) + OH-(aq) + CuBr2(s) ® NaBr2(aq) + CuOH(s) C. Na+(aq) + Br-(aq) ® NaBr(s) D. 2Na+(aq) + Br2(aq) ® 2NaBr(s) E. 2OH-(aq) + Cu2+(aq) ® Cu(OH)2(s)
14. What is the net ionic equation for the reaction of aqueous sodium hydroxide and aqueous iron(II) chloride? A. Na+(aq) + OH–(aq) ® NaOH(s) B. Na+(aq) + Cl–(aq) ® NaCl(s) C. Fe2+(aq) + 2OH–(aq) ® Fe(OH)2(s) D. Fe2+(aq) + OH–(aq) ® FeOH+(s) E. Fe2+(aq) + 2Cl–(aq) ® FeCl2(s)
15. What net ionic reaction occurs when aqueous solutions of potassium carbonate and iron(III) bromide are mixed? A. CO32-(aq) + 2Fe+(aq) ® Fe2CO3(s) B. 3CO32-(aq) + 2Fe3+(aq) ® Fe2(CO3)3(s) C. 3CO32-(aq) + 6Fe+(aq) ® 3Fe2CO3(s) D. 3K2CO3(aq) + 2Fe3+(aq) + 6Br-(aq) ® Fe2(CO3)3(s) + 6KBr(s) E. no reaction occurs
16. What reaction occurs when aqueous solutions of silver nitrate and potassium iodide are mixed? A. Ag+(aq) + K+(aq) ® AgK(s) B. NO3-(aq) + I-(aq) ® NO3I(s) C. NO3-(aq) + K+(aq) ® KNO3(s) D. Ag+(aq) + NO3-(aq) + K+(aq) + I-(aq) ® AgI(s) + 2KNO3(s) E. Ag+(aq) + I-(aq) ® AgI(s)
17. Which of the following reagents can be added to silver nitrate to precipitate the silver ion? A. Pb(NO3)2 B. NH4ClO4 C. KNO3 D. NaI E. KClO4
18. Which of the following reagents can be added to a solution of sodium carbonate to precipitate the carbonate ion? A. FeCl3 B. MgBr2 C. Cu(ClO4)2 D. Pb(NO3)2 E. all of the above
19. Write a balanced net ionic equation for the reaction of manganese(II) hydroxide, Mn(OH)2, with hydrochloric acid. A. Mn(OH)2(s) + 2H+(aq) ® Mn2+(aq) + 2H2O(l) B. Mn(OH)2(s) + 2H+(aq) ® MnH2(s) + 2OH-(aq) C. (OH)3(aq) + 3H+(aq) ® 3H2O(l) D. 3OH-(aq) + 3H+(aq) ® 3H2O(l) E. Mn2+(aq) + 2Cl-(aq) ® MnCl2(aq)
20. Identify all of the spectator ions in the reaction below. Zn(OH)2(s) + 2H+(aq) + 2NO3-(aq) ® Zn2+(aq) + 2NO3-(aq) + 2H2O(l) A. Zn(OH)2 B. NO3C. Zn2+ D. H+ E. H+ and NO3-
21. Write a balanced net ionic equation for the reaction of aqueous solutions of lead(II) nitrate and potassium chloride. A. Pb(NO3)2(aq) + 2KCl(aq) ® PbCl2(s) + 2KNO3(aq) B. Pb2+(aq) + 2K+(aq) ® PbK2(s) C. Pb2+(aq) + 2Cl-(aq) ® PbCl2(s) D. NO3-(aq) + K+(aq) ® KNO3(s) E. no precipitation occurs.
22. Identify all of the spectator ions in the precipitation reaction below. Ca2+(aq) + 2Br-(aq) + 2Li+(aq) + CO32-(aq) ® CaCO3(s) + 2Li+(aq) + 2Br-(aq) A. Ca2+ and Li+ B. Br- and CO32C. Br- and Li+ D. CaCO3 E. Ca2+, Br-, Li+, and CO32-
23. What volume of 0.250 M KCl(aq) will completely react with 50.0 mL of 0.115 M Pb(NO3)2(aq)? Pb2+(aq) + 2Cl-(aq) ® PbCl2(s) A. 23.0 mL B. 46.0 mL C. 11.5 mL D. 109 mL E. 218 mL
24. What mass of precipitate is formed when 25.00 mL of 0.200 M BaCl2 and 35.00 mL of 0.125 M Na2SO4 are mixed? A. 0.510 g B. 1.17 g C. 1.02 g D. 2.04 g E. 0.583 g
25. According to the Arrhenius acid-base definition A. acids produce H+ in aqueous solutions and bases produce OH- in aqueous solutions. B. acids produce OH- in aqueous solutions and bases produce H+ in aqueous solutions. C. acids only react with bases. D. all hydrogen halides are strong acids. E. acids and bases are strong electrolytes.
26. Which list contains only strong acids? A. HCl, HNO3, HF, HClO4 B. H2SO4, H3PO4, HClO4, NH3 C. HCl, HNO3, H3PO4, HClO4 D. HCl, H2SO4, HClO4, HI E. HNO3, H2SO4, NaOH, H3PO4
27. Which of the following compounds is a weak acid? A. HCl B. H3PO4 C. HNO3 D. HClO4 E. H2SO4
28. Which of the following are strong bases: NH3, NaOH, Ba(OH)2, and HF? A. NH3 and HF B. NaOH and Ba(OH)2 C. NH3 and NaOH D. NaOH, Ba(OH)2, and HF E. NH3, NaOH, Ba(OH)2, and HF
29. Which of the following compounds is a weak base? A. NaOH B. H2CO3 C. LiCl D. NH3 E. CH3CO2H
30. What is the net ionic equation for the reaction of aqueous hydrochloric acid and aqueous potassium hydroxide? A. HCl(aq) + OH–(aq) ® H2O(l) + Cl–(aq) B. Cl–(aq) + K+(aq) ® KCl(s) C. HCl(aq) + KOH(aq) ® KCl(aq) + H2O(l) D. Cl–(aq) + K+(aq) ® KCl(aq) E. H+(aq) + OH–(aq) ® H2O(l)
31. What is the balanced net ionic equation for the reaction between aqueous solutions of acetic acid and sodium hydroxide? A. CH3CO2H(aq) + OH-(aq) ® CH3(aq) + CO32-(aq) + H2(g) B. CH3CO2H(aq) + 3NaOH(aq) ® CCO2H3-(aq) + 3Na+(aq) + H2O(l) C. CH3CO2H(aq) + OH-(aq) ® CH3-(aq) + CO2(g) + H2O(l) D. CH3CO2H(aq) + OH-(aq) ® CH3CO2-(aq) + H2O(l) E. CH3CO2H(aq) + 2Na+(aq) + 2OH-(aq) ® CH3CO22-(aq) + 2Na+(aq) + H2O(l)
32. When HCl(g) and NH3(g) are mixed, a white solid forms. What is the balanced equation for this reaction? A. HCl(g) + NH3(g) ® NH4Cl(s) B. HCl(g) + NH3(g) ® NH2Cl(g) + H2(s) C. HCl(g) + NH3(g) ® NH4(s) + Cl(g) D. HCl(g) + NH3(g) ® NH2Cl(s) + H2(g) E. 3HCl(g) + NH3(g) ® NCl3(s) + 3H2(g)
33. What are the spectator ions in the reaction between aqueous hydrobromic acid and aqueous sodium hydroxide? A. Na+ only B. H+ and OH– C. Na+ and Br– D. Br– only E. H+, Br–, Na+, and OH–
34. In the laboratory, acid spills are often neutralized by adding sodium bicarbonate. What mass of sodium bicarbonate reacts with 225 mL of 6.00 M HCl? H+(aq) + NaHCO3(s) ® H2O(l) + CO2(g) + Na+(aq) A. 1.35 g B. 71.5 g C. 113 g D. 143 g E. 2240 g
35. Exactly 19.36 mL of 0.1481 M NaOH is used to titrate a 25.00 mL sample of HF. What is the concentration of the hydrofluoric acid? A. 6.464 ´ 10-2 M B. 1.147 ´ 10-1 M C. 1.916 ´ 10-1 M D. 2.867 ´ 10-3 M E. 6.570 ´ 10-3 M
36. A 25.00 mL sample of vinegar is diluted with water to a final volume of 250.0 mL. A 25.00 mL portion of the diluted vinegar is then titrated with 0.0998 M NaOH. If 22.43 mL of NaOH is required to reach the equivalence point, what is the original concentration of acetic acid in the vinegar? CH3CO2H(aq) + OH-(aq) ® CH3CO2-(aq) + H2O(l) A. 0.0538 M B. 0.0895 M C. 0.538 M D. 0.895 M E. 5.38 M
37. A mass of 0.4113 g of an unknown acid, HA, is titrated with NaOH(aq). If the acid reacts with 28.10 mL of 0.1055 M NaOH(aq), what is the molar mass of the acid? A. 2.965 ´ 10–3 g/mol B. 9.128 g/mol C. 138.7 g/mol D. 337.3 g/mol E. 820.7 g/mol
38. The principal ingredient in Tums antacid tablets is calcium carbonate, CaCO3. A single tablet contains 0.500 g CaCO3. What volume of 0.2500 M HCl is required to titrate a tablet of Tums? 2H+(aq) + CO32-(aq) ® H2O(l) + CO2(g) A. 0.0100 mL B. 0.0200 mL C. 10.0 mL D. 20.0 mL E. 40.0 mL
39. A oxidizing agent is a species that A. takes a proton from an Arrhenius acid. B. is oxidized in a chemical reaction. C. gains electrons in a chemical reaction. D. loses electrons in a chemical reaction. E. gives a proton to an Arrhenius base.
40. Assign oxidation numbers to each atom in sodium hydrogen carbonate, NaHCO3? A. Na = +1, H = –1, C = +6, O = –2 B. Na = +1, H = +1, C = +4, O = –2 C. Na = +1, H = –1, C = +2, O = –2 D. Na = –1, H = +1, C = 0, O = –2 E. Na = 0, H = 0, C = 0, O = 0
41. Assign oxidation numbers to each atom in manganese(IV) oxide. A. Mn = +4; O = 0 B. Mn = +4; O = -2 C. Mn = +2; O = 0 D. Mn = +2; O = -2 E. Mn = 0; O = 0
42. What is the oxidation number of iodine in potassium iodate, KIO3? A. –1 B. 0 C. +3 D. +5 E. +7
43. Identify the oxidizing and reducing agents in the redox reaction below. 3Cu(s) + 8H+(aq) + 2NO3-(aq) ® 3Cu2+(aq) + 2NO(g) + 4H2O(l) A. reducing agent: Cu; oxidizing agent: Cu2+ B. reducing agent: NO3-; oxidizing agent: NO C. reducing agent: Cu; oxidizing agent: NO3D. reducing agent: NO3-; oxidizing agent: Cu E. reducing agent: H+; oxidizing agent: NO
44. Balance the reaction and identify which species is reduced and which is oxidized. Na(s) + H2O(l) ® Na+(aq) + OH-(aq) + H2(g) A. Na(s) + H2O(l) ® Na+(aq) + OH-(aq) + H2(g); Na reduced, H2O oxidized B. 2Na(s) + 2H2O(l) ® 2Na+(aq) + 2OH-(aq) + H2(g); Na reduced, H2O oxidized C. Na(s) + H2O(l) ® Na+(aq) + OH-(aq) + H2(g); H2O reduced, Na oxidized D. 2Na(s) + 2H2O(l) ® 2Na+(aq) + 2OH-(aq) + H2(g); H2O reduced, Na oxidized E. 2Na(s) + 2H2O(l) ® 2Na+(aq) + 2OH-(aq) + H2(g); not a redox reaction
45. Which of the following chemical equations show oxidation-reduction reactions? 1. 2. 3.
Mg(s) + I2(aq) ® MgI2(s) Pb(ClO4)2(aq) + 2KI(aq) ® PbI2(s) + 2KClO4(aq) Fe2O3(s) + 3CO(g) ® 2 Fe(s) + 3CO2(g)
A. 1 only B. 2 only C. 1 and 2 D. 1 and 3 E. 2 and 3 46. What is the oxidation half-reaction for the reaction of zinc with hydrochloric acid? Zn(s) + 2H+(aq) ® Zn2+(aq) + H2(g) A. Zn(s) ® Zn2+(aq) + 2eB. Zn(s) + 2e- ® Zn2+(aq) C. 2H+(aq) ® H2(g) + 2eD. 2H+(aq) + 2e- ® H2(g) E. none of the above
47. What is the reduction half-reaction for the equation below? 5Fe2+(aq) + MnO4-(aq) + 8H+(aq) ® 5Fe3+(aq) + Mn2+(aq) + 4H2O(l) A. MnO4-(aq) + 5e- ® Mn2+(aq) + 2O2(g) B. MnO4-(aq) + 8H+(aq) + 5e- ® Mn2+(aq) + 4H2O(l) C. Fe2+(aq) + e- ® Fe3+(aq) D. 8H+(aq) + 8e- ® 8H2O(l) E. none of the above
48. Write a balanced half-reaction for the reduction of permanganate ion, MnO7-, to Mn2+ in an acidic solution. A. MnO4-(aq) + 4H2O(l) ® Mn2+(aq) + 8OH-(aq) B. MnO4-(aq) + 8H+(aq) + 5e- ® Mn2+(aq) + 4H2O(l) C. MnO4-(aq) + 5e- ® Mn2+(aq) + 4O2-(aq) D. MnO4-(aq) ® Mn2+(aq) + 4O2-(aq) E. MnO4-(aq) + 4H+(aq) + 5e- ® Mn2+(aq) + 4OH-(aq)
49. Identify the reaction type for the reaction between potassium metal and hydrochloric acid. 2K(s) + 2H+(aq) ® 2K+(aq) + H2(g) A. precipitation B. acid-base C. oxidation-reduction D. both answers a and c E. none of the above
50. Which of the pictoral representations best represents the precipitation reaction that occurs between aqueous solutions of Fe3+ and OH-? Assume that the circles represent cations and the squares represent anions.
A.
B.
C.
D.
E.
51. What volume of 0.2500 M cobalt(III) sulfate is required to react completely with 5.00 g of sodium carbonate? A. 126 mL B. 786 mL C. 102 mL D. 117 mL E. 314 mL
52. In the reaction given below, how many grams of sodium metal are consumed if 2.02 g of hydrogen gas are produced? 2Na(s) + 2H2O(l) ® 2NaOH(aq) + H2(g) A. 92.0 g B. 5.75 g C. 11.5 g D. 23.0 g E. 46.0 g
53. How many grams of KClO3 are needed to produce of 2.56 grams of O2? 2KClO3(s) ® 2KCl(s) + 3O2(g) A. 1.00 g B. 1.71 g C. 6.53 g D. 9.80 g E. 14.7 g
54. Polythiophene is a polymer with unique properties. What property is might make it useful in uniforms for soldiers? A. Electrotension, the fibers are super-strong under electric charge. B. Thermochromism, the fibers change color with temperature changes. C. Polymorphism, the fibers change shape under stress. D. Electrochromism, the fibers change colors under electric charge. E. none of the above
Chapter 4--Reactions in Aqueous Solution Key
1. Which of the following statements concerning aqueous NaCl is/are CORRECT? 1. 2. 3.
Sodium chloride is a strong electrolyte. Sodium chloride is insoluble in water. NaCl(aq) consists of neutral NaCl molecules dissolved in water.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 2. Which of the following pairs of aqueous solutions will produce a precipitate when mixed? A. NaCl(aq) and Ca(NO3)2(aq) B. KCl(aq) and NaOH(aq) C. Cu(NO3)2(aq) and AgNO3(aq) D. ZnSO4(aq) and LiCl(aq) E. K2SO4(aq) and Pb(ClO4)2(aq)
3. Which of the following pairs of aqueous solutions will NOT produce a precipitate when mixed? A. AgNO3(aq) and KI(aq) B. Cu(NO3)2(aq) and MgCl2(aq) C. AlCl3(aq) and KOH(aq) D. BaBr2(aq) and Na2SO4(aq) E. AgNO3(aq) and NaCl(aq)
4. Precipitation reactions occur A. when group 1 cations are mixed with group 17 anions. B. when insoluble reactants are mixed. C. when ionic compounds react to form non-ionic products. D. predominantly with halide salts. E. when soluble ionic reactants combine to form insoluble products.
5. Which of the following statements is/are correct? 1. 2. 3.
Most ionic compounds containing nitrate ion are soluble in water. Most ionic compounds containing sulfate ion are insoluble in water. Most ionic compounds containing carbonate ion are soluble in water.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 6. Which of the following statements is/are correct? 1. 2. 3.
Most ionic compounds containing phosphate ion are insoluble in water. Most ionic compounds containing potassium ion are insoluble in water. Most ionic compounds containing hydroxide ion are soluble in water.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 7. If aqueous solutions of sodium sulfate and barium chloride are mixed, a white precipitate forms. What is the identity of the precipitate? A. Na2Ba B. NaCl2 C. NaCl D. BaSO4 E. Ba2SO4
8. If aqueous solutions of nickel(II) chloride and potassium phosphate are mixed, which insoluble compound will form? A. Ni3(PO4)2 B. NiCl2 C. KCl D. K2Ni E. K3PO4
9. Which of the following ionic compounds is not soluble in water? A. K2CO3 B. AgNO3 C. CaBr2 D. CsI E. PbI2
10. Which of the following ionic compounds are likely to be soluble in water: Mg(OH)2, Pb(NO3)2, AgI, Na2CO3, and Cu3(PO4)3? A. Na2CO3 only B. Pb(NO3)2 and Na2CO3 C. Mg(OH)2, Na2CO3, and Cu3(PO4)3 D. Mg(OH)2, AgI, and Cu3(PO4)3 E. Pb(NO3)2, AgI, Na2CO3, and Cu3(PO4)3
11. Which of the following compounds is least soluble in water? A. Fe(OH)3. B. Sr(OH)2. C. Al(NO3)3. D. KI. E. CuCl2.
12. If an aqueous solution of ____ is added to a mixture of Pb2+ and Ba2+, the lead ion will precipitate, but the barium ion will remain in solution. A. NaOH B. Na2SO4 C. K3PO4 D. KCO3 E. Ca(CH3CO2)2
13. What reaction occurs when aqueous solutions of sodium hydroxide and copper(II) bromide are mixed? A. Na+(aq) + OH-(aq) + Cu2+(aq) + 2Br-(aq) ® NaCu(s) + Br2OH(aq) B. Na+(aq) + OH-(aq) + CuBr2(s) ® NaBr2(aq) + CuOH(s) C. Na+(aq) + Br-(aq) ® NaBr(s) D. 2Na+(aq) + Br2(aq) ® 2NaBr(s) E. 2OH-(aq) + Cu2+(aq) ® Cu(OH)2(s)
14. What is the net ionic equation for the reaction of aqueous sodium hydroxide and aqueous iron(II) chloride? A. Na+(aq) + OH–(aq) ® NaOH(s) B. Na+(aq) + Cl–(aq) ® NaCl(s) C. Fe2+(aq) + 2OH–(aq) ® Fe(OH)2(s) D. Fe2+(aq) + OH–(aq) ® FeOH+(s) E. Fe2+(aq) + 2Cl–(aq) ® FeCl2(s)
15. What net ionic reaction occurs when aqueous solutions of potassium carbonate and iron(III) bromide are mixed? A. CO32-(aq) + 2Fe+(aq) ® Fe2CO3(s) B. 3CO32-(aq) + 2Fe3+(aq) ® Fe2(CO3)3(s) C. 3CO32-(aq) + 6Fe+(aq) ® 3Fe2CO3(s) D. 3K2CO3(aq) + 2Fe3+(aq) + 6Br-(aq) ® Fe2(CO3)3(s) + 6KBr(s) E. no reaction occurs
16. What reaction occurs when aqueous solutions of silver nitrate and potassium iodide are mixed? A. Ag+(aq) + K+(aq) ® AgK(s) B. NO3-(aq) + I-(aq) ® NO3I(s) C. NO3-(aq) + K+(aq) ® KNO3(s) D. Ag+(aq) + NO3-(aq) + K+(aq) + I-(aq) ® AgI(s) + 2KNO3(s) E. Ag+(aq) + I-(aq) ® AgI(s)
17. Which of the following reagents can be added to silver nitrate to precipitate the silver ion? A. Pb(NO3)2 B. NH4ClO4 C. KNO3 D. NaI E. KClO4
18. Which of the following reagents can be added to a solution of sodium carbonate to precipitate the carbonate ion? A. FeCl3 B. MgBr2 C. Cu(ClO4)2 D. Pb(NO3)2 E. all of the above
19. Write a balanced net ionic equation for the reaction of manganese(II) hydroxide, Mn(OH)2, with hydrochloric acid. A. Mn(OH)2(s) + 2H+(aq) ® Mn2+(aq) + 2H2O(l) B. Mn(OH)2(s) + 2H+(aq) ® MnH2(s) + 2OH-(aq) C. (OH)3(aq) + 3H+(aq) ® 3H2O(l) D. 3OH-(aq) + 3H+(aq) ® 3H2O(l) E. Mn2+(aq) + 2Cl-(aq) ® MnCl2(aq)
20. Identify all of the spectator ions in the reaction below. Zn(OH)2(s) + 2H+(aq) + 2NO3-(aq) ® Zn2+(aq) + 2NO3-(aq) + 2H2O(l) A. Zn(OH)2 B. NO3C. Zn2+ D. H+ E. H+ and NO3-
21. Write a balanced net ionic equation for the reaction of aqueous solutions of lead(II) nitrate and potassium chloride. A. Pb(NO3)2(aq) + 2KCl(aq) ® PbCl2(s) + 2KNO3(aq) B. Pb2+(aq) + 2K+(aq) ® PbK2(s) C. Pb2+(aq) + 2Cl-(aq) ® PbCl2(s) D. NO3-(aq) + K+(aq) ® KNO3(s) E. no precipitation occurs.
22. Identify all of the spectator ions in the precipitation reaction below. Ca2+(aq) + 2Br-(aq) + 2Li+(aq) + CO32-(aq) ® CaCO3(s) + 2Li+(aq) + 2Br-(aq) A. Ca2+ and Li+ B. Br- and CO32C. Br- and Li+ D. CaCO3 E. Ca2+, Br-, Li+, and CO32-
23. What volume of 0.250 M KCl(aq) will completely react with 50.0 mL of 0.115 M Pb(NO3)2(aq)? Pb2+(aq) + 2Cl-(aq) ® PbCl2(s) A. 23.0 mL B. 46.0 mL C. 11.5 mL D. 109 mL E. 218 mL
24. What mass of precipitate is formed when 25.00 mL of 0.200 M BaCl2 and 35.00 mL of 0.125 M Na2SO4 are mixed? A. 0.510 g B. 1.17 g C. 1.02 g D. 2.04 g E. 0.583 g
25. According to the Arrhenius acid-base definition A. acids produce H+ in aqueous solutions and bases produce OH- in aqueous solutions. B. acids produce OH- in aqueous solutions and bases produce H+ in aqueous solutions. C. acids only react with bases. D. all hydrogen halides are strong acids. E. acids and bases are strong electrolytes.
26. Which list contains only strong acids? A. HCl, HNO3, HF, HClO4 B. H2SO4, H3PO4, HClO4, NH3 C. HCl, HNO3, H3PO4, HClO4 D. HCl, H2SO4, HClO4, HI E. HNO3, H2SO4, NaOH, H3PO4
27. Which of the following compounds is a weak acid? A. HCl B. H3PO4 C. HNO3 D. HClO4 E. H2SO4
28. Which of the following are strong bases: NH3, NaOH, Ba(OH)2, and HF? A. NH3 and HF B. NaOH and Ba(OH)2 C. NH3 and NaOH D. NaOH, Ba(OH)2, and HF E. NH3, NaOH, Ba(OH)2, and HF
29. Which of the following compounds is a weak base? A. NaOH B. H2CO3 C. LiCl D. NH3 E. CH3CO2H
30. What is the net ionic equation for the reaction of aqueous hydrochloric acid and aqueous potassium hydroxide? A. HCl(aq) + OH–(aq) ® H2O(l) + Cl–(aq) B. Cl–(aq) + K+(aq) ® KCl(s) C. HCl(aq) + KOH(aq) ® KCl(aq) + H2O(l) D. Cl–(aq) + K+(aq) ® KCl(aq) E. H+(aq) + OH–(aq) ® H2O(l)
31. What is the balanced net ionic equation for the reaction between aqueous solutions of acetic acid and sodium hydroxide? A. CH3CO2H(aq) + OH-(aq) ® CH3(aq) + CO32-(aq) + H2(g) B. CH3CO2H(aq) + 3NaOH(aq) ® CCO2H3-(aq) + 3Na+(aq) + H2O(l) C. CH3CO2H(aq) + OH-(aq) ® CH3-(aq) + CO2(g) + H2O(l) D. CH3CO2H(aq) + OH-(aq) ® CH3CO2-(aq) + H2O(l) E. CH3CO2H(aq) + 2Na+(aq) + 2OH-(aq) ® CH3CO22-(aq) + 2Na+(aq) + H2O(l)
32. When HCl(g) and NH3(g) are mixed, a white solid forms. What is the balanced equation for this reaction? A. HCl(g) + NH3(g) ® NH4Cl(s) B. HCl(g) + NH3(g) ® NH2Cl(g) + H2(s) C. HCl(g) + NH3(g) ® NH4(s) + Cl(g) D. HCl(g) + NH3(g) ® NH2Cl(s) + H2(g) E. 3HCl(g) + NH3(g) ® NCl3(s) + 3H2(g)
33. What are the spectator ions in the reaction between aqueous hydrobromic acid and aqueous sodium hydroxide? A. Na+ only B. H+ and OH– C. Na+ and Br– D. Br– only E. H+, Br–, Na+, and OH–
34. In the laboratory, acid spills are often neutralized by adding sodium bicarbonate. What mass of sodium bicarbonate reacts with 225 mL of 6.00 M HCl? H+(aq) + NaHCO3(s) ® H2O(l) + CO2(g) + Na+(aq) A. 1.35 g B. 71.5 g C. 113 g D. 143 g E. 2240 g
35. Exactly 19.36 mL of 0.1481 M NaOH is used to titrate a 25.00 mL sample of HF. What is the concentration of the hydrofluoric acid? A. 6.464 ´ 10-2 M B. 1.147 ´ 10-1 M C. 1.916 ´ 10-1 M D. 2.867 ´ 10-3 M E. 6.570 ´ 10-3 M
36. A 25.00 mL sample of vinegar is diluted with water to a final volume of 250.0 mL. A 25.00 mL portion of the diluted vinegar is then titrated with 0.0998 M NaOH. If 22.43 mL of NaOH is required to reach the equivalence point, what is the original concentration of acetic acid in the vinegar? CH3CO2H(aq) + OH-(aq) ® CH3CO2-(aq) + H2O(l) A. 0.0538 M B. 0.0895 M C. 0.538 M D. 0.895 M E. 5.38 M
37. A mass of 0.4113 g of an unknown acid, HA, is titrated with NaOH(aq). If the acid reacts with 28.10 mL of 0.1055 M NaOH(aq), what is the molar mass of the acid? A. 2.965 ´ 10–3 g/mol B. 9.128 g/mol C. 138.7 g/mol D. 337.3 g/mol E. 820.7 g/mol
38. The principal ingredient in Tums antacid tablets is calcium carbonate, CaCO3. A single tablet contains 0.500 g CaCO3. What volume of 0.2500 M HCl is required to titrate a tablet of Tums? 2H+(aq) + CO32-(aq) ® H2O(l) + CO2(g) A. 0.0100 mL B. 0.0200 mL C. 10.0 mL D. 20.0 mL E. 40.0 mL
39. A oxidizing agent is a species that A. takes a proton from an Arrhenius acid. B. is oxidized in a chemical reaction. C. gains electrons in a chemical reaction. D. loses electrons in a chemical reaction. E. gives a proton to an Arrhenius base.
40. Assign oxidation numbers to each atom in sodium hydrogen carbonate, NaHCO3? A. Na = +1, H = –1, C = +6, O = –2 B. Na = +1, H = +1, C = +4, O = –2 C. Na = +1, H = –1, C = +2, O = –2 D. Na = –1, H = +1, C = 0, O = –2 E. Na = 0, H = 0, C = 0, O = 0
41. Assign oxidation numbers to each atom in manganese(IV) oxide. A. Mn = +4; O = 0 B. Mn = +4; O = -2 C. Mn = +2; O = 0 D. Mn = +2; O = -2 E. Mn = 0; O = 0
42. What is the oxidation number of iodine in potassium iodate, KIO3? A. –1 B. 0 C. +3 D. +5 E. +7
43. Identify the oxidizing and reducing agents in the redox reaction below. 3Cu(s) + 8H+(aq) + 2NO3-(aq) ® 3Cu2+(aq) + 2NO(g) + 4H2O(l) A. reducing agent: Cu; oxidizing agent: Cu2+ B. reducing agent: NO3-; oxidizing agent: NO C. reducing agent: Cu; oxidizing agent: NO3D. reducing agent: NO3-; oxidizing agent: Cu E. reducing agent: H+; oxidizing agent: NO
44. Balance the reaction and identify which species is reduced and which is oxidized. Na(s) + H2O(l) ® Na+(aq) + OH-(aq) + H2(g) A. Na(s) + H2O(l) ® Na+(aq) + OH-(aq) + H2(g); Na reduced, H2O oxidized B. 2Na(s) + 2H2O(l) ® 2Na+(aq) + 2OH-(aq) + H2(g); Na reduced, H2O oxidized C. Na(s) + H2O(l) ® Na+(aq) + OH-(aq) + H2(g); H2O reduced, Na oxidized D. 2Na(s) + 2H2O(l) ® 2Na+(aq) + 2OH-(aq) + H2(g); H2O reduced, Na oxidized E. 2Na(s) + 2H2O(l) ® 2Na+(aq) + 2OH-(aq) + H2(g); not a redox reaction
45. Which of the following chemical equations show oxidation-reduction reactions? 1. 2. 3.
Mg(s) + I2(aq) ® MgI2(s) Pb(ClO4)2(aq) + 2KI(aq) ® PbI2(s) + 2KClO4(aq) Fe2O3(s) + 3CO(g) ® 2 Fe(s) + 3CO2(g)
A. 1 only B. 2 only C. 1 and 2 D. 1 and 3 E. 2 and 3 46. What is the oxidation half-reaction for the reaction of zinc with hydrochloric acid? Zn(s) + 2H+(aq) ® Zn2+(aq) + H2(g) A. Zn(s) ® Zn2+(aq) + 2eB. Zn(s) + 2e- ® Zn2+(aq) C. 2H+(aq) ® H2(g) + 2eD. 2H+(aq) + 2e- ® H2(g) E. none of the above
47. What is the reduction half-reaction for the equation below? 5Fe2+(aq) + MnO4-(aq) + 8H+(aq) ® 5Fe3+(aq) + Mn2+(aq) + 4H2O(l) A. MnO4-(aq) + 5e- ® Mn2+(aq) + 2O2(g) B. MnO4-(aq) + 8H+(aq) + 5e- ® Mn2+(aq) + 4H2O(l) C. Fe2+(aq) + e- ® Fe3+(aq) D. 8H+(aq) + 8e- ® 8H2O(l) E. none of the above
48. Write a balanced half-reaction for the reduction of permanganate ion, MnO7-, to Mn2+ in an acidic solution. A. MnO4-(aq) + 4H2O(l) ® Mn2+(aq) + 8OH-(aq) B. MnO4-(aq) + 8H+(aq) + 5e- ® Mn2+(aq) + 4H2O(l) C. MnO4-(aq) + 5e- ® Mn2+(aq) + 4O2-(aq) D. MnO4-(aq) ® Mn2+(aq) + 4O2-(aq) E. MnO4-(aq) + 4H+(aq) + 5e- ® Mn2+(aq) + 4OH-(aq)
49. Identify the reaction type for the reaction between potassium metal and hydrochloric acid. 2K(s) + 2H+(aq) ® 2K+(aq) + H2(g) A. precipitation B. acid-base C. oxidation-reduction D. both answers a and c E. none of the above
50. Which of the pictoral representations best represents the precipitation reaction that occurs between aqueous solutions of Fe3+ and OH-? Assume that the circles represent cations and the squares represent anions.
A.
B.
C.
D.
E.
51. What volume of 0.2500 M cobalt(III) sulfate is required to react completely with 5.00 g of sodium carbonate? A. 126 mL B. 786 mL C. 102 mL D. 117 mL E. 314 mL
52. In the reaction given below, how many grams of sodium metal are consumed if 2.02 g of hydrogen gas are produced? 2Na(s) + 2H2O(l) ® 2NaOH(aq) + H2(g) A. 92.0 g B. 5.75 g C. 11.5 g D. 23.0 g E. 46.0 g
53. How many grams of KClO3 are needed to produce of 2.56 grams of O2? 2KClO3(s) ® 2KCl(s) + 3O2(g) A. 1.00 g B. 1.71 g C. 6.53 g D. 9.80 g E. 14.7 g
54. Polythiophene is a polymer with unique properties. What property is might make it useful in uniforms for soldiers? A. Electrotension, the fibers are super-strong under electric charge. B. Thermochromism, the fibers change color with temperature changes. C. Polymorphism, the fibers change shape under stress. D. Electrochromism, the fibers change colors under electric charge. E. none of the above
Chapter 5--Gases 1. All of the following are units of gas pressure EXCEPT A. the newton. B. the standard atmosphere. C. the bar. D. millimeters of mercury. E. the pascal.
2. Convert 618 mm Hg to kPa. (1 atm = 760 mm Hg = 101.3 kPa) A. 0.0121 kPa B. 1.23 kPa C. 0.813 kPa D. 82.4 kPa E. 4.64 ´ 103 kPa
3. A balloon with a volume of 8.73 L contains 0.321 moles of helium gas. What is the density of the gas? A. 0.0368 g/L B. 0.147 g/L C. 0.700 g/L D. 2.80 g/L E. 27.1 g/L
4. If the volume of a confined gas is expanded to four times the original volume while its temperature remains constant, what change will be observed? A. The pressure of the gas will decrease to 1/4 its original value. B. The pressure of the gas will decrease to 1/2 its original value. C. The pressure of the gas will remain unchanged. D. The pressure of the gas will increase to twice its original value. E. The pressure of the gas will increase to four times its original value.
5. Place the following units of pressure in order from lowest to highest pressure. A. 1 atm < 1 Pa < 1 mm Hg < 1 bar B. 1 mm Hg < 1 bar < 1 atm < 1 Pa C. 1 Pa < 1 mm Hg < 1 bar < 1 atm D. 1 Pa < 1 mm Hg < 1 atm < 1 bar E. 1 bar < 1 mm Hg < 1 Pa < 1 atm
6. All of the following relationships are false for gases EXCEPT A. volume is inversely proportional to the moles of gas. B. volume is directly proportional to pressure in mm Hg. C. volume is directly proportional to pressure in atmospheres. D. volume is directly proportional to temperature in Kelvin. E. volume is directly proportional to the gas constant R.
7. Avogadro's law states that A. 1 liter of any gas contains 6.02 ´ 1023 gas molecules. B. the volume of a gas is directly proportional to its temperature. C. the gas constant equals 0.0821 Latm/(molK) for all ideal gases. D. the volume of a gas must always be a constant. E. equal volumes of all gases at the same pressure and temperature contain an equal number of moles.
8. The molar masses of helium and oxygen are 4.0 g/mol and 16 g/mol, respectively. At the same temperature and pressure, 1 mole of helium will occupy A. the same volume as 1 mole of oxygen. B. four times the volume of 1 mole of oxygen. C. twice the volume of 1 mole of oxygen. D. half the volume of 1 mole of oxygen. E. one fourth the volume of 1 mole of oxygen.
9. A balloon is filled with He gas to a volume of 2.10 L at 35C. The balloon is placed in liquid nitrogen until its temperature reaches –196C. Assuming the pressure remains constant, what is the volume of the cooled balloon? A. -0.375 L B. 0.375 L C. 0.525 L D. 0.00909 L E. 8.40 L
10. A tightly sealed 4.0-L flask contains 884 mm Hg of N2 at 94.0C. The flask is cooled until the pressure is reduced 442 mm Hg. What is the temperature of the gas? A. –47.0C B. 47.0C C. –89.5C D. 184C E. 188C
11. A bicycle tire is filled to a pressure of 4.42 atm (65 psi) at a temperature of 12C. If the temperature of the tire increases to 33C, what is the pressure in the tire? Assume the volume of the tire is constant. A. 4.75 atm B. 5.91 atm C. 9.28 atm D. 12.1 atm E. 16.1 atm
12. If 8.5 g of oxygen gas (O2) is introduced into an evacuated 1.50 L flask at 22C, what is the pressure inside the flask? A. 0.32 atm B. 0.48 atm C. 4.3 atm D. 6.4 atm E. 137 atm
13. At 108C, the pressure in a 10.0 L flask is 874 mm Hg. How many moles of gas are in the flask? A. 0.368 mol B. 0.873 mol C. 1.30 mol D. 348 mol E. 986 mol
14. What volume does 22.4 moles of hydrogen gas occupy at 0C and 1.00 atm? A. 0.0821 L B. 1.00 L C. 22.4 L D. 184 L E. 502 L
15. A 3.00 L flask contains 2.33 g of argon gas at 312 mm Hg. What is the temperature of the gas? A. 151 K B. 257 K C. 292 K D. 341 K E. 4890 K
16. Which of the following relationships is/are CORRECT for gases? 1. 2. 3.
The amount of a gas (in moles) is inversely proportional to its volume (at constant temperature and pressure). The volume of a gas is directly proportional to its temperature in kelvin (at constant pressure and moles). The pressure of a gas is inversely proportional to its temperature in kelvin (at constant volume and moles).
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 2 and 3 17. A mass of 1.37 g of an unknown gas is introduced into an evacuated 1.70 L flask. If the pressure in the flask is 0.814 atm at 98C, which of the following gases might be in the flask? (R = 0.0821 Latm/molK) A. CH4 B. C2H2 C. C2H6 D. C3H8 E. C4H10
18. A mass of 5.0 grams of dry ice, CO2(s), is sealed in an evacuated 2.0 L plastic soda bottle. What is the pressure inside the bottle when the CO2 is heated to 35C? A. 0.16 atm B. 1.4 atm C. 8.9 atm D. 18 atm E. 63 atm
19. At 28C, a gas cylinder containing hydrogen has an internal volume and pressure of 46.6 L and 1.50 ´ 102 atm. The gas is allowed to escape the cylinder and expand until its pressure reaches 1.00 atmosphere. If the temperature of the gas remains constant, what volume will the gas occupy? A. 0.311 L B. 2.22 L C. 46.6 L D. 6.99 ´ 103 L E. 8.91 ´ 103 L
20. At 21C, a gas cylinder containing nitrogen has an internal volume and pressure of 46.6 L and 1.50 ´ 102 atm. What mass of nitrogen gas is contained in the cylinder? A. 2.90 ´ 102 g B. 4.05 ´ 103 g C. 8.11 ´ 103 g D. 1.14 ´ 105 g E. none of the above
21. At 338 mm Hg and 72C, a sample of carbon monoxide gas occupies a volume of 0.225 L. The gas is transferred to a 1.50-L flask and the temperature is reduced to –15C. What is the pressure of the gas in the flask? A. 8.91 mm Hg B. 37.9 mm Hg C. 67.8 mm Hg D. 3.018 ´ 103 mm Hg E. 5.48 ´ 104 mm Hg
22. What volume of Ar at 45C and 1.25 atm contains the same number of particles as 0.655 L of H2 at 15C and 2.00 atm? A. 0.452 L B. 1.16 L C. 0.949 L D. 0.371 L E. 3.14 L
23. A balloon is filled with 1.50 L of helium gas at sea level, 1.00 atm and 32C. The balloon is released and it rises to an altitude of 30,000 ft. If the pressure at this altitude is 228 mm Hg and the temperature is -45C, what is the volume of the balloon? A. 0.336 L B. 1.56 L C. 1.68 L D. 2.81 L E. 3.74 L
24. A volatile compound with a mass of 0.8822 grams is placed in an evacuated 0.250 L flask. The flask is heated to evaporate the compound. At 99C, the pressure in the flask is 1.25 atm. What is the molar mass of the compound? A. 86.2 g/mol B. 116 g/mol C. 229 g/mol D. 257 g/mol E. 303 g/mol
25. Calculate the density (in g/L) of CH4(g) at 75C and 2.1 atm. (R = 0.08206 Latm/molK) A. 0.18 g/L B. 0.85 g/L C. 1.2 g/L D. 3.2 g/L E. 5.5 g/L
26. At 1.00 km above sea level, the atmospheric pressure is 675 mm Hg and the temperature is 282 K. If nitrogen comprises 78.1% (mole percent) of air, what is the density of nitrogen at this height? A. 0.0238 g/L B. 0.521 g/L C. 0.839 g/L D. 1.07 g/L E. 638 g/L
27. At what temperature does 1.00 atm of carbon dioxide (CO2) gas have the same density as 1.00 atm of helium gas at 25C? A. 0 K B. 25 K C. 482 K D. 983 K E. 3.28 ´ 103 K
28. Which of the following gases has the greatest density at 35 C and 450 mm Hg? A. CH4 B. Ar C. N2 D. Cl2 E. C3H8
29. The density of hydrogen gas in a flask is 0.147 g/L at 305 K. What is the pressure inside the flask? A. 0.139 atm B. 1.19 atm C. 1.83 atm D. 2.98 atm E. 3.69 atm
30. An unknown gas contains 85.6% C and 14.4% H. At 0.455 atm and 425 K, the gas has a density of 0.549 g/L. What is the molecular formula for the gas? A. CH4 B. C2H4 C. C3H6 D. C4H6 E. C4H8
31. If 6.46 L of gaseous ethanol reacts with 16.1 L O2, what is the maximum volume of gaseous carbon dioxde produced? Assume that the temperature of the reactants and products is 425C and the pressure remains constant at 1.00 atm. CH3CH2OH(g) + 3O2(g) ® 2CO2(g) + 3H2O(g) A. 6.46 L B. 10.7 L C. 12.9 L D. 16.1 L E. 22.6 L
32. A volume of 3.0 L of butane is burned in excess oxygen. Balance the chemical equation below and determine how many total liters of gases are produced. Assume that both the reactant and product temperature is 500 K and the pressure of the system remains constant at 1.0 atm. C4H10(g) + O2(g) ® CO2(g) + H2O(g) A. 3.0 L B. 6.0 L C. 12 L D. 27 L E. 42 L
33. If 5.00 L of propane is burned in 21.0 L of oxygen, what volume of carbon dioxide is produced? Assume that the temperature of the reactants and products is 25C and the pressure of the system remains constant at 1.0 atm. C3H8(g) + 5O2(g) ® 3CO2(g) + 4H2O(l) A. 12.6 L B. 15.0 L C. 21.0 L D. 25.6 L E. 26.0 L
34. A 10.0 L flask contains 2.5 atm of ethane gas and 8.0 atm of oxygen gas at 28C. The contents of the flask react until the limiting reactant is consumed. What is the pressure of carbon dioxide in the flask after the temperature returns to 28C? 2C2H6(g) + 7O2(g) ® 4CO2(g) + 6H2O(g) A. 2.5 atm B. 4.6 atm C. 5.0 atm D. 5.5 atm E. 10.5 atm
35. What volume of O2, measured at 225C and 0.970 atm, will be produced by the decomposition of 3.16 g KClO3? (R = 0.08206 Latm/molK) 2KClO3(s) ® 2KCl(s) + 3O2(g) A. 1.09 L B. 1.24 L C. 1.63 L D. 3.26 L E. 52.1 L
36. Aqueous hydrochloric acid reacts with magnesium to produce hydrogen gas according to the balanced equation below. 2H+(aq) + Mg(s) ® Mg2+(aq) + H2(g) If 250.0 mL of 3.00 M HCl is combined with 9.92 g Mg, what volume of hydrogen gas can be produced? Assume the temperature and pressure of the gas are 25C and 0.988 atm, respectively. (R = 0.0821 Latm/molK) A. 4.60 L B. 5.05 L C. 9.29 L D. 10.1 L E. 18.6 L
37. Ammonia gas is produced commercially from the reaction of nitrogen and hydrogen. What volume of ammonia can be produced from the reaction of 5.5 ´ 103 kg of N2 and 1.5 ´ 103 kg of H2? Assume the reaction is 100% efficient and the product is collected at 325 K and 25 atm. N2(g) + 3H2(g) ® 2NH3(g) A. 7.0 ´ 103 L B. 1.4 ´ 104 L C. 1.8 ´ 104 L D. 4.2 ´ 105 L E. 5.2 ´ 105 L
38. Nitroglycerin (227.1 g/mol) releases a large amount of energy and gaseous products upon decomposition. If 10.00 grams of nitroglycerin decomposes in an evacuated 1.00 L flask, what is the pressure inside the flask? Assume the temperature is 4.00 ´ 102 K and that the flask survives the explosion. 4C3H5N3O9(l) ® 6N2(g) + 12CO2(g) + 10H2O(g) + O2(g) A. 1.45 atm B. 10.5 atm C. 41.9 atm D. 328 atm E. 952 atm
39. The composition (in mole percent) of the atmosphere is 78.1% N2, 21.0% O2, and 0.9% Ar. What is the partial pressure of each gas when the barometric pressure is 754.1 mm Hg? A. N2 = 21.1 atm, O2 = 6.5 atm, Ar = 0.3 atm B. N2 = 78.1 mm Hg, O2 = 21.0 mm Hg, Ar = 0.9 mm Hg C. N2 = 244 mm Hg, O2 = 244 mm Hg, Ar = 244 mm Hg D. N2 = 405 mm Hg, O2 = 234 mm Hg, Ar = 293 mm Hg E. N2 = 589 mm Hg, O2 = 158 mm Hg, Ar = 7 mm Hg
40. Sulfur burns in oxygen with a deep blue flame to produce sulfur dioxide. If 5.85 g S8 and 1.00 atm of O2 completely react in a 5.00 L flask (at 25C), determine the partial pressure of SO2 (at 25C) and the total pressure in the flask. S8(s) + 8O2(g) ® 8SO2(g) A. O2 = 0 atm, SO2 = 1.00 atm, total pressure = 1.00 atm B. O2 = 0.093 atm, SO2 = 0.917 atm, total pressure = 1.00 atm C. O2 = 0.107 atm, SO2 = 0.893 atm, total pressure = 1.00 atm D. O2 = 0.855 atm, SO2 = 0.145 atm, total pressure = 1.00 atm E. O2 = 0.917 atm, SO2 = 0.163 atm, total pressure = 1.08 atm
41. An unknown mass of ammonium perchlorate, NH4ClO4 (117.5 g/mol), is placed in an evacuated 1.00 L flask and heated to 251C. At this temperature the NH4ClO4 decomposes violently. The gaseous products exert a pressure of 466 mm Hg at 251C. What mass of NH4ClO4 was placed in the flask? 2NH4ClO4(s) ® N2(g) + Cl2(g) + 2O2(g) + 4H2O(g) A. 0.149 g B. 0.419 g C. 0.682 g D. 1.67 g E. 3.19 g
42. Water can be decomposed by electrolysis to hydrogen gas and oxygen gas. What mass of water must decompose to yield 24.0 L of oxygen gas at 1.00 atm and 25C? 2H2O(l) ® 2H2(g) + O2(g) A. 11.1 g B. 17.7 g C. 23.6 g D. 35.3 g E. 70.7 g
43. A mixture of H2 and O2 is placed in a 5.00 L flask at 22 C. The partial pressure of the H2 is 2.7 atm and the partial pressure of the O2 is 1.5 atm. What is the mole fraction of H2? A. 0.13 B. 0.36 C. 0.56 D. 0.64 E. 0.87
44. A sample of carbon dioxide is collected over water at 25C (vapor pressure H2O(l) = 23.8 mm Hg). The CO2 and water vapor occupy a volume of 1.80 L at a pressure of 783.0 mm Hg. What mass of CO2 is present? A. 3.23 g B. 4.40 g C. 9.02 g D. 14.7 g E. 16.3 g
45. A 10.0 L flask is used to collect 0.500 moles of N2 and 0.180 moles of O2 over water at 30C. What is the pressure in the flask? (vapor pressure H2O(l) = 31.8 mm Hg) A. -30.1 atm B. 1.15 atm C. 1.48 atm D. 1.69 atm E. 1.73 atm
46. Water can be decomposed by electrolysis to hydrogen gas and oxygen gas. If 2.33 g of water is decomposed to H2(g) and O2(g) and the gases are collected in a 1.00 L flask over water at 25C (vapor pressure H2O(l) = 23.8 mm Hg), what is the pressure in the flask? A. 3.19 atm B. 4.71 atm C. 4.75 atm D. 4.78 atm E. 9.52 atm
47. Which of the following statements are postulates of the kinetic-molecular theory of gases? 1. 2. 3.
Gas particles are in constant, random motion. Collisions between gas molecules are elastic. Gas pressure is caused by collisions of gas molecules with the walls of the container.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 48. At a given temperature, molecules of different gases A. have the same average kinetic energy of translational motion. B. have the same average velocity. C. have the same diameter. D. have the same density. E. have identical masses.
49. The average speed of a gas molecule is A. inversely proportional to the square root of its mass. B. inversely proportional to the gas constant, R. C. directly proportional to the square of its temperature in K. D. inversely proportional to its kinetic energy. E. directly proportional to the square of its temperature in C.
50. At 25C, the average speed of a hydrogen molecule is 1.92 ´ 103 m/s. What is the average speed of a nitrogen molecule at the same temperature? A. 138 m/s B. 515 m/s C. 7.16 ´ 103 m/s D. 2.67 ´ 104 m/s E. none of the above
51. SF6(g) can be used as an insulating gas between glass panes of a window. If the temperature of the gas is 10C, what is the average speed of the gas? A. 41.3 m/s B. 372 m/s C. 706 m/s D. 484 m/s E. 220 m/s
52. At what temperature does O2 have the same average speed as H2 does at 273 K? A. 273 K B. 389 K C. 498 K D. 658 K E. 4.33 ´ 103 K
53. Methane gas, CH4, effuses through a barrier at a rate of 0.568 mL/minute. If an unknown gas effuses through the same barrier at a rate of 0.343 mL/minute, what is the molar mass of the gas? A. 20.8 g/mol B. 28.0 g/mol C. 32.0 g/mol D. 44.0 g/mol E. 64.0 g/mol
54. In an experiment, argon is allowed to effuse through a tiny opening into an evacuated 5.00 ´ 102 mL flask for 30.0 seconds, at which point the pressure in the flask is found to be 15.0 mm Hg. The experiment is repeated using an unknown gas at the same temperature and pressure. After 30.0 seconds, the pressure is found to be 47.4 mm Hg. What is the molar mass of the gas? A. 4.00 g/mol B. 16.0 g/mol C. 28.0 g/mol D. 32.0 g/mol E. 83.8 g/mol
55. Which conditions will cause the greatest deviation from the ideal gas law? A. 100 atm and 500 K B. 100 atm and 10 K C. 0.001 atm and 500 K D. 0.001 and 10 K E. 0.001 and 273 K
56. One way in which real gases differ from ideal gases is that the molecules of a real gas A. have no kinetic energy. B. move in curved paths. C. have no mass. D. occupy no volume. E. are attracted to each other.
57. Which of the following statements concerning real gases is/are CORRECT? 1. 2. 3.
Real gases are always liquids or solids at temperatures below 273.15 K. The pressure of a real gas is higher than predicted by the ideal gas law. The molecules in a real gas are attracted to each other.
A. 1 only B. 2 only C. 3 only D. 2 and 3 E. 1, 2 and 3 58. Sodium azide is used to generate the gas that fills the airbag in an automobile crash. What is the gas generated? A. Na (g) B. N2 (g) C. N3 (g) D. Az (g) E. O2 (g)
Chapter 5--Gases Key
1. All of the following are units of gas pressure EXCEPT A. the newton. B. the standard atmosphere. C. the bar. D. millimeters of mercury. E. the pascal.
2. Convert 618 mm Hg to kPa. (1 atm = 760 mm Hg = 101.3 kPa) A. 0.0121 kPa B. 1.23 kPa C. 0.813 kPa D. 82.4 kPa E. 4.64 ´ 103 kPa
3. A balloon with a volume of 8.73 L contains 0.321 moles of helium gas. What is the density of the gas? A. 0.0368 g/L B. 0.147 g/L C. 0.700 g/L D. 2.80 g/L E. 27.1 g/L
4. If the volume of a confined gas is expanded to four times the original volume while its temperature remains constant, what change will be observed? A. The pressure of the gas will decrease to 1/4 its original value. B. The pressure of the gas will decrease to 1/2 its original value. C. The pressure of the gas will remain unchanged. D. The pressure of the gas will increase to twice its original value. E. The pressure of the gas will increase to four times its original value.
5. Place the following units of pressure in order from lowest to highest pressure. A. 1 atm < 1 Pa < 1 mm Hg < 1 bar B. 1 mm Hg < 1 bar < 1 atm < 1 Pa C. 1 Pa < 1 mm Hg < 1 bar < 1 atm D. 1 Pa < 1 mm Hg < 1 atm < 1 bar E. 1 bar < 1 mm Hg < 1 Pa < 1 atm
6. All of the following relationships are false for gases EXCEPT A. volume is inversely proportional to the moles of gas. B. volume is directly proportional to pressure in mm Hg. C. volume is directly proportional to pressure in atmospheres. D. volume is directly proportional to temperature in Kelvin. E. volume is directly proportional to the gas constant R.
7. Avogadro's law states that A. 1 liter of any gas contains 6.02 ´ 1023 gas molecules. B. the volume of a gas is directly proportional to its temperature. C. the gas constant equals 0.0821 Latm/(molK) for all ideal gases. D. the volume of a gas must always be a constant. E. equal volumes of all gases at the same pressure and temperature contain an equal number of moles.
8. The molar masses of helium and oxygen are 4.0 g/mol and 16 g/mol, respectively. At the same temperature and pressure, 1 mole of helium will occupy A. the same volume as 1 mole of oxygen. B. four times the volume of 1 mole of oxygen. C. twice the volume of 1 mole of oxygen. D. half the volume of 1 mole of oxygen. E. one fourth the volume of 1 mole of oxygen.
9. A balloon is filled with He gas to a volume of 2.10 L at 35C. The balloon is placed in liquid nitrogen until its temperature reaches –196C. Assuming the pressure remains constant, what is the volume of the cooled balloon? A. -0.375 L B. 0.375 L C. 0.525 L D. 0.00909 L E. 8.40 L
10. A tightly sealed 4.0-L flask contains 884 mm Hg of N2 at 94.0C. The flask is cooled until the pressure is reduced 442 mm Hg. What is the temperature of the gas? A. –47.0C B. 47.0C C. –89.5C D. 184C E. 188C
11. A bicycle tire is filled to a pressure of 4.42 atm (65 psi) at a temperature of 12C. If the temperature of the tire increases to 33C, what is the pressure in the tire? Assume the volume of the tire is constant. A. 4.75 atm B. 5.91 atm C. 9.28 atm D. 12.1 atm E. 16.1 atm
12. If 8.5 g of oxygen gas (O2) is introduced into an evacuated 1.50 L flask at 22C, what is the pressure inside the flask? A. 0.32 atm B. 0.48 atm C. 4.3 atm D. 6.4 atm E. 137 atm
13. At 108C, the pressure in a 10.0 L flask is 874 mm Hg. How many moles of gas are in the flask? A. 0.368 mol B. 0.873 mol C. 1.30 mol D. 348 mol E. 986 mol
14. What volume does 22.4 moles of hydrogen gas occupy at 0C and 1.00 atm? A. 0.0821 L B. 1.00 L C. 22.4 L D. 184 L E. 502 L
15. A 3.00 L flask contains 2.33 g of argon gas at 312 mm Hg. What is the temperature of the gas? A. 151 K B. 257 K C. 292 K D. 341 K E. 4890 K
16. Which of the following relationships is/are CORRECT for gases? 1. 2. 3.
The amount of a gas (in moles) is inversely proportional to its volume (at constant temperature and pressure). The volume of a gas is directly proportional to its temperature in kelvin (at constant pressure and moles). The pressure of a gas is inversely proportional to its temperature in kelvin (at constant volume and moles).
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 2 and 3 17. A mass of 1.37 g of an unknown gas is introduced into an evacuated 1.70 L flask. If the pressure in the flask is 0.814 atm at 98C, which of the following gases might be in the flask? (R = 0.0821 Latm/molK) A. CH4 B. C2H2 C. C2H6 D. C3H8 E. C4H10
18. A mass of 5.0 grams of dry ice, CO2(s), is sealed in an evacuated 2.0 L plastic soda bottle. What is the pressure inside the bottle when the CO2 is heated to 35C? A. 0.16 atm B. 1.4 atm C. 8.9 atm D. 18 atm E. 63 atm
19. At 28C, a gas cylinder containing hydrogen has an internal volume and pressure of 46.6 L and 1.50 ´ 102 atm. The gas is allowed to escape the cylinder and expand until its pressure reaches 1.00 atmosphere. If the temperature of the gas remains constant, what volume will the gas occupy? A. 0.311 L B. 2.22 L C. 46.6 L D. 6.99 ´ 103 L E. 8.91 ´ 103 L
20. At 21C, a gas cylinder containing nitrogen has an internal volume and pressure of 46.6 L and 1.50 ´ 102 atm. What mass of nitrogen gas is contained in the cylinder? A. 2.90 ´ 102 g B. 4.05 ´ 103 g C. 8.11 ´ 103 g D. 1.14 ´ 105 g E. none of the above
21. At 338 mm Hg and 72C, a sample of carbon monoxide gas occupies a volume of 0.225 L. The gas is transferred to a 1.50-L flask and the temperature is reduced to –15C. What is the pressure of the gas in the flask? A. 8.91 mm Hg B. 37.9 mm Hg C. 67.8 mm Hg D. 3.018 ´ 103 mm Hg E. 5.48 ´ 104 mm Hg
22. What volume of Ar at 45C and 1.25 atm contains the same number of particles as 0.655 L of H2 at 15C and 2.00 atm? A. 0.452 L B. 1.16 L C. 0.949 L D. 0.371 L E. 3.14 L
23. A balloon is filled with 1.50 L of helium gas at sea level, 1.00 atm and 32C. The balloon is released and it rises to an altitude of 30,000 ft. If the pressure at this altitude is 228 mm Hg and the temperature is -45C, what is the volume of the balloon? A. 0.336 L B. 1.56 L C. 1.68 L D. 2.81 L E. 3.74 L
24. A volatile compound with a mass of 0.8822 grams is placed in an evacuated 0.250 L flask. The flask is heated to evaporate the compound. At 99C, the pressure in the flask is 1.25 atm. What is the molar mass of the compound? A. 86.2 g/mol B. 116 g/mol C. 229 g/mol D. 257 g/mol E. 303 g/mol
25. Calculate the density (in g/L) of CH4(g) at 75C and 2.1 atm. (R = 0.08206 Latm/molK) A. 0.18 g/L B. 0.85 g/L C. 1.2 g/L D. 3.2 g/L E. 5.5 g/L
26. At 1.00 km above sea level, the atmospheric pressure is 675 mm Hg and the temperature is 282 K. If nitrogen comprises 78.1% (mole percent) of air, what is the density of nitrogen at this height? A. 0.0238 g/L B. 0.521 g/L C. 0.839 g/L D. 1.07 g/L E. 638 g/L
27. At what temperature does 1.00 atm of carbon dioxide (CO2) gas have the same density as 1.00 atm of helium gas at 25C? A. 0 K B. 25 K C. 482 K D. 983 K E. 3.28 ´ 103 K
28. Which of the following gases has the greatest density at 35 C and 450 mm Hg? A. CH4 B. Ar C. N2 D. Cl2 E. C3H8
29. The density of hydrogen gas in a flask is 0.147 g/L at 305 K. What is the pressure inside the flask? A. 0.139 atm B. 1.19 atm C. 1.83 atm D. 2.98 atm E. 3.69 atm
30. An unknown gas contains 85.6% C and 14.4% H. At 0.455 atm and 425 K, the gas has a density of 0.549 g/L. What is the molecular formula for the gas? A. CH4 B. C2H4 C. C3H6 D. C4H6 E. C4H8
31. If 6.46 L of gaseous ethanol reacts with 16.1 L O2, what is the maximum volume of gaseous carbon dioxde produced? Assume that the temperature of the reactants and products is 425C and the pressure remains constant at 1.00 atm. CH3CH2OH(g) + 3O2(g) ® 2CO2(g) + 3H2O(g) A. 6.46 L B. 10.7 L C. 12.9 L D. 16.1 L E. 22.6 L
32. A volume of 3.0 L of butane is burned in excess oxygen. Balance the chemical equation below and determine how many total liters of gases are produced. Assume that both the reactant and product temperature is 500 K and the pressure of the system remains constant at 1.0 atm. C4H10(g) + O2(g) ® CO2(g) + H2O(g) A. 3.0 L B. 6.0 L C. 12 L D. 27 L E. 42 L
33. If 5.00 L of propane is burned in 21.0 L of oxygen, what volume of carbon dioxide is produced? Assume that the temperature of the reactants and products is 25C and the pressure of the system remains constant at 1.0 atm. C3H8(g) + 5O2(g) ® 3CO2(g) + 4H2O(l) A. 12.6 L B. 15.0 L C. 21.0 L D. 25.6 L E. 26.0 L
34. A 10.0 L flask contains 2.5 atm of ethane gas and 8.0 atm of oxygen gas at 28C. The contents of the flask react until the limiting reactant is consumed. What is the pressure of carbon dioxide in the flask after the temperature returns to 28C? 2C2H6(g) + 7O2(g) ® 4CO2(g) + 6H2O(g) A. 2.5 atm B. 4.6 atm C. 5.0 atm D. 5.5 atm E. 10.5 atm
35. What volume of O2, measured at 225C and 0.970 atm, will be produced by the decomposition of 3.16 g KClO3? (R = 0.08206 Latm/molK) 2KClO3(s) ® 2KCl(s) + 3O2(g) A. 1.09 L B. 1.24 L C. 1.63 L D. 3.26 L E. 52.1 L
36. Aqueous hydrochloric acid reacts with magnesium to produce hydrogen gas according to the balanced equation below. 2H+(aq) + Mg(s) ® Mg2+(aq) + H2(g) If 250.0 mL of 3.00 M HCl is combined with 9.92 g Mg, what volume of hydrogen gas can be produced? Assume the temperature and pressure of the gas are 25C and 0.988 atm, respectively. (R = 0.0821 Latm/molK) A. 4.60 L B. 5.05 L C. 9.29 L D. 10.1 L E. 18.6 L
37. Ammonia gas is produced commercially from the reaction of nitrogen and hydrogen. What volume of ammonia can be produced from the reaction of 5.5 ´ 103 kg of N2 and 1.5 ´ 103 kg of H2? Assume the reaction is 100% efficient and the product is collected at 325 K and 25 atm. N2(g) + 3H2(g) ® 2NH3(g) A. 7.0 ´ 103 L B. 1.4 ´ 104 L C. 1.8 ´ 104 L D. 4.2 ´ 105 L E. 5.2 ´ 105 L
38. Nitroglycerin (227.1 g/mol) releases a large amount of energy and gaseous products upon decomposition. If 10.00 grams of nitroglycerin decomposes in an evacuated 1.00 L flask, what is the pressure inside the flask? Assume the temperature is 4.00 ´ 102 K and that the flask survives the explosion. 4C3H5N3O9(l) ® 6N2(g) + 12CO2(g) + 10H2O(g) + O2(g) A. 1.45 atm B. 10.5 atm C. 41.9 atm D. 328 atm E. 952 atm
39. The composition (in mole percent) of the atmosphere is 78.1% N2, 21.0% O2, and 0.9% Ar. What is the partial pressure of each gas when the barometric pressure is 754.1 mm Hg? A. N2 = 21.1 atm, O2 = 6.5 atm, Ar = 0.3 atm B. N2 = 78.1 mm Hg, O2 = 21.0 mm Hg, Ar = 0.9 mm Hg C. N2 = 244 mm Hg, O2 = 244 mm Hg, Ar = 244 mm Hg D. N2 = 405 mm Hg, O2 = 234 mm Hg, Ar = 293 mm Hg E. N2 = 589 mm Hg, O2 = 158 mm Hg, Ar = 7 mm Hg
40. Sulfur burns in oxygen with a deep blue flame to produce sulfur dioxide. If 5.85 g S8 and 1.00 atm of O2 completely react in a 5.00 L flask (at 25C), determine the partial pressure of SO2 (at 25C) and the total pressure in the flask. S8(s) + 8O2(g) ® 8SO2(g) A. O2 = 0 atm, SO2 = 1.00 atm, total pressure = 1.00 atm B. O2 = 0.093 atm, SO2 = 0.917 atm, total pressure = 1.00 atm C. O2 = 0.107 atm, SO2 = 0.893 atm, total pressure = 1.00 atm D. O2 = 0.855 atm, SO2 = 0.145 atm, total pressure = 1.00 atm E. O2 = 0.917 atm, SO2 = 0.163 atm, total pressure = 1.08 atm
41. An unknown mass of ammonium perchlorate, NH4ClO4 (117.5 g/mol), is placed in an evacuated 1.00 L flask and heated to 251C. At this temperature the NH4ClO4 decomposes violently. The gaseous products exert a pressure of 466 mm Hg at 251C. What mass of NH4ClO4 was placed in the flask? 2NH4ClO4(s) ® N2(g) + Cl2(g) + 2O2(g) + 4H2O(g) A. 0.149 g B. 0.419 g C. 0.682 g D. 1.67 g E. 3.19 g
42. Water can be decomposed by electrolysis to hydrogen gas and oxygen gas. What mass of water must decompose to yield 24.0 L of oxygen gas at 1.00 atm and 25C? 2H2O(l) ® 2H2(g) + O2(g) A. 11.1 g B. 17.7 g C. 23.6 g D. 35.3 g E. 70.7 g
43. A mixture of H2 and O2 is placed in a 5.00 L flask at 22 C. The partial pressure of the H2 is 2.7 atm and the partial pressure of the O2 is 1.5 atm. What is the mole fraction of H2? A. 0.13 B. 0.36 C. 0.56 D. 0.64 E. 0.87
44. A sample of carbon dioxide is collected over water at 25C (vapor pressure H2O(l) = 23.8 mm Hg). The CO2 and water vapor occupy a volume of 1.80 L at a pressure of 783.0 mm Hg. What mass of CO2 is present? A. 3.23 g B. 4.40 g C. 9.02 g D. 14.7 g E. 16.3 g
45. A 10.0 L flask is used to collect 0.500 moles of N2 and 0.180 moles of O2 over water at 30C. What is the pressure in the flask? (vapor pressure H2O(l) = 31.8 mm Hg) A. -30.1 atm B. 1.15 atm C. 1.48 atm D. 1.69 atm E. 1.73 atm
46. Water can be decomposed by electrolysis to hydrogen gas and oxygen gas. If 2.33 g of water is decomposed to H2(g) and O2(g) and the gases are collected in a 1.00 L flask over water at 25C (vapor pressure H2O(l) = 23.8 mm Hg), what is the pressure in the flask? A. 3.19 atm B. 4.71 atm C. 4.75 atm D. 4.78 atm E. 9.52 atm
47. Which of the following statements are postulates of the kinetic-molecular theory of gases? 1. 2. 3.
Gas particles are in constant, random motion. Collisions between gas molecules are elastic. Gas pressure is caused by collisions of gas molecules with the walls of the container.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 48. At a given temperature, molecules of different gases A. have the same average kinetic energy of translational motion. B. have the same average velocity. C. have the same diameter. D. have the same density. E. have identical masses.
49. The average speed of a gas molecule is A. inversely proportional to the square root of its mass. B. inversely proportional to the gas constant, R. C. directly proportional to the square of its temperature in K. D. inversely proportional to its kinetic energy. E. directly proportional to the square of its temperature in C.
50. At 25C, the average speed of a hydrogen molecule is 1.92 ´ 103 m/s. What is the average speed of a nitrogen molecule at the same temperature? A. 138 m/s B. 515 m/s C. 7.16 ´ 103 m/s D. 2.67 ´ 104 m/s E. none of the above
51. SF6(g) can be used as an insulating gas between glass panes of a window. If the temperature of the gas is 10C, what is the average speed of the gas? A. 41.3 m/s B. 372 m/s C. 706 m/s D. 484 m/s E. 220 m/s
52. At what temperature does O2 have the same average speed as H2 does at 273 K? A. 273 K B. 389 K C. 498 K D. 658 K E. 4.33 ´ 103 K
53. Methane gas, CH4, effuses through a barrier at a rate of 0.568 mL/minute. If an unknown gas effuses through the same barrier at a rate of 0.343 mL/minute, what is the molar mass of the gas? A. 20.8 g/mol B. 28.0 g/mol C. 32.0 g/mol D. 44.0 g/mol E. 64.0 g/mol
54. In an experiment, argon is allowed to effuse through a tiny opening into an evacuated 5.00 ´ 102 mL flask for 30.0 seconds, at which point the pressure in the flask is found to be 15.0 mm Hg. The experiment is repeated using an unknown gas at the same temperature and pressure. After 30.0 seconds, the pressure is found to be 47.4 mm Hg. What is the molar mass of the gas? A. 4.00 g/mol B. 16.0 g/mol C. 28.0 g/mol D. 32.0 g/mol E. 83.8 g/mol
55. Which conditions will cause the greatest deviation from the ideal gas law? A. 100 atm and 500 K B. 100 atm and 10 K C. 0.001 atm and 500 K D. 0.001 and 10 K E. 0.001 and 273 K
56. One way in which real gases differ from ideal gases is that the molecules of a real gas A. have no kinetic energy. B. move in curved paths. C. have no mass. D. occupy no volume. E. are attracted to each other.
57. Which of the following statements concerning real gases is/are CORRECT? 1. 2. 3.
Real gases are always liquids or solids at temperatures below 273.15 K. The pressure of a real gas is higher than predicted by the ideal gas law. The molecules in a real gas are attracted to each other.
A. 1 only B. 2 only C. 3 only D. 2 and 3 E. 1, 2 and 3 58. Sodium azide is used to generate the gas that fills the airbag in an automobile crash. What is the gas generated? A. Na (g) B. N2 (g) C. N3 (g) D. Az (g) E. O2 (g)
Chapter 6--Electronic Structure and the Periodic Table 1. The Navy uses electromagnetic radiation of extremely long wavelengths to communicate with submerged submarines. The Navy's ELF (Extremely Low Frequency) systems emit 76 s-1 radiation from facilities located in Wisconsin and Michigan. What is the wavelength of this radiation (in meters)? A. 2.1 ´ 10-7 m B. 1.3 ´ 10-3 m C. 2.5 ´ 103 m D. 3.9 ´ 106 m E. 2.3 ´ 1010 m
2. If a radio wave has a frequency of 17.0 MHz, what is the wavelength of this radiation? A. 5.10 ´ 1015 m B. 5.67 ´ 10–2 m C. 5.88 ´ 10–2 m D. 1.76 ´ 101 m E. 1.76 ´ 106 m
3. If a cordless phone operates at a frequency of 9.00 ´ 108 s–1. What is the wavelength of this radiation (in nm)? A. 0.333 m B. 5.96 ´ 10–25 m C. 3.33 ´ 10–9 m D. 3.71 ´ 10–18 m E. 2.70 ´ 1017 m
4. What is the frequency of a gamma ray radiation that has a wavelength of 11.4 pm? A. 3.80 ´ 10–20 s–1 B. 1.74 ´ 10–14 s–1 C. 2.63 ´ 107 s–1 D. 3.42 ´ 109 s–1 E. 2.63 ´ 1019 s–1
5. Green laser pointers emit radiation at 532 nm. What is the frequency of this radiation? A. 8.12 ´ 1013 Hz B. 5.64 ´ 1014 Hz C. 1.60 ´ 1015 Hz D. 9.10 ´ 1015 Hz E. 1.60 ´ 1016 Hz
6. Which of the following regions of the electromagnetic spectrum have longer wavelengths than visible light? 1. 2. 3.
infrared radiation ultraviolet radiation microwave radiation
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1 and 3 7. Place the following regions of the electromagnetic spectrum in order from highest to lowest energy. A. infrared > gamma rays > x-rays > visible > radio B. gamma rays > x-rays > visible > infrared > radio C. x-rays > gamma rays > infrared > visible > radio D. radio > x-rays > gamma rays > visible > infrared E. visible > infrared > radio > x-rays > gamma rays
8. The ____ of a photon of light is ____ proportional to its frequency and ____ proportional to its wavelength. A. energy, directly, inversely B. energy, inversely, directly C. velocity, directly, inversely D. intensity, inversely, directly E. amplitude, directly, inversely
9. Some digital cordless phones operate at 2.4 GHz (1 GHz = 109 Hz). What is the energy, in joules, of a single 2.4 GHz photon? A. 8.3 ´ 10-35 J B. 1.6 ´ 10-24 J C. 5.4 ´ 10-19 J D. 0.96 J E. 0.13 J
10. Excited sodium atoms emit light in the infrared at 589 nm. What is the energy of a single photon with this wavelength? A. 5.09 ´ 1014 J B. 1.12 ´ 10–27 J C. 3.37 ´ 10–19 J D. 3.37 ´ 10–28 J E. 1.30 ´ 10–-19 J
11. When Cs-137 decays, it emits gamma radiation. The energy of one photon is 1.06 ´ 10-13 J. What is the wavelength of this radiation? A. 1.87 ´ 10-12 m B. 2.08 ´ 10-12 m C. 2.44 ´ 10-12 m D. 1.88 ´ 10-11 m E. 1.58 ´ 10-10 m
12. If the energy of 1.00 mole of photons is 245 kJ, what is the wavelength of the light? A. 122 nm B. 488 nm C. 1220 nm D. 787 nm E. 811 nm
13. According to the Bohr model for the hydrogen atom, the energy necessary to excite an electron from n = 2 to n = 3 is ____ the energy necessary to excite an electron from n = 3 to n = 4. A. greater than B. equal to C. less than D. either equal to or greater than E. either less than or equal to
14. Calculate the wavelength of the line in the Lyman series that results from the transition n = 3 to n = 1. The Rydberg constant equals 2.180 ´ 10-18 J. A. 45.59 nm B. 72.81 nm C. 91.12 nm D. 102.5 nm E. 136.7 nm
15. Calculate the energy of a photon in the Balmer series that results from the transition n = 3 to n = 2. What is the region of the electromagnetic spectrum? The Rydberg constant equals 2.180 ´ 10-18 J. A. 3.028 ´ 10-19 J; visible B. 4.578 ´ 10-19 J; infrared C. 3.633 ´ 10-19 J; visible D. 3.633 ´ 10-18 J; ultraviolet E. 2.180 ´ 10-18 J; ultraviolet
16. A line in the Pfund series (nlo = 5) occurs at 3.74 ´ 10-6 m. What is nhi for this transition? The Rydberg constant equals 2.180 ´ 10-18 J. A. 6 B. 7 C. 8 D. 9 E. 10
17. For which of the following transitions would a hydrogen atom absorb a photon with the longest wavelength? A. n = 1 to n = 2 B. n = 4 to n = 6 C. n = 5 to n = 4 D. n = 7 to n = 6 E. n = 6 to n = 7
18. The Schrödinger wave equation A. proves electrons have positive and negative spins. B. calculates the precise position and momentum of an electron at any given time. C. is used to compute the wavelength of small particles. D. can be solved to find the probability of finding an electron in a region of space. E. proves that photons are particles.
19. The Pauli exclusion principle states that A. no two electrons from a given atom can have the same spin. B. no two electrons from a given atom can have the same four quantum numbers. C. two electrons can occupy an orbital if they have the same spin. D. two electrons can occupy an orbital if they have opposite spins. E. two electrons can occupy an orbital if they have opposite charges.
20. What type of orbital is designated n = 3, A. 3s B. 3p C. 3d D. 2f E. 2d
= 2,
= –2?
21. Which of the following sets of quantum numbers refers to a 2s orbital? A. n = 1,
= 2,
= 2, ms = +
B. n = 1,
= 2,
= 1, ms = +
C. n = 2,
= 2,
= 0, ms = +
D. n = 2,
= 1,
= -1, ms = +
E. n = 2,
= 0,
= 0, ms = +
22. All of the following sets of quantum numbers are allowed EXCEPT A. n = 1,
= 0,
= 1, ms = -
.
B. n = 2,
= 1,
= 0, ms = +
.
C. n = 3,
= 1,
= -1, ms = -
.
D. n = 4,
= 3,
= -1, ms = -
.
E. n = 6,
= 3,
= -3, ms = +
.
23. What is the total number of orbitals having n = 4 and A. 3 B. 5 C. 7 D. 9 E. 10
= 2?
24. How many orbitals have the following set of quantum numbers: n = 5, A. 0 B. 1 C. 3 D. 6 E. 7
25. What is the total capacity of electrons in n = 3? A. 5 B. 10 C. 18 D. 32 E. 50
26. What is the total capacity of electrons in n = 5, A. 2 B. 6 C. 10 D. 14 E. 32
= 3?
27. Which of the following sets of quantum numbers refers to a 4f orbital? A. n = 4,
= 1,
= 0, ms = +
B. n = 4,
= 2,
= -1, ms = -
C. n = 4,
= 3,
= -2, ms = -
D. n = 4,
= 4,
= 0, ms = +
E. n = 4,
= 4,
= -4, ms = +
28. What type of orbital is designated by n = 5, A. 5d B. 2f C. 5p D. 2s E. 5p
= 2,
= +1?
= 3,
= +2?
29. Which of the following properties is associated with the value of the A. the shape of an orbital B. the size of an orbital C. the number of electrons in an orbital D. the energy of an orbital E. the orientation in space of an orbital
30. What is the ground state electron configuration of 25Mn? A. 1s22s22p63s23p63d64s1 B. 1s22s22p63s23p63d54s2 C. 1s22s22p63s23p63d7 D. 1s22s22p63s23p63d5 E. 1s22s22p63s23p64s24d5
31. What is the electron configuration of 80Hg? A. [Xe]6s26d10 B. [Kr]5d106s2 C. [Xe]4f145d106s2 D. [Xe]5d106s2 E. [Kr]4d104f146s2
32. What is the ground state electron configuration of 94Pu? A. [Xe]5f66s2 B. [Xe]6f67s2 C. [Rn]6f67s2 D. [Rn]5f67s2 E. [Rn]5f66d107s2
33. What is the electron configuration of S2-? A. 1s22s22p6 B. 1s22s22p63s2 C. 1s22s22p63s23p4 D. 1s22s22p63s23p6 E. 1s22s22p63s23p64s1
quantum number?
34. What is the ground state electron configuration for Cr3+? A. [Ar] B. [Ar]3d74s2 C. [Ar]3d14s2 D. [Ar]3d24s1 E. [Ar]3d3
35. What is the electron configuration of Ga+? A. [Ar]3d104s2 B. [Ar]3d104p2 C. [Ar]3d84s24p2 D. [Ar]3d104s24d4 E. [Ar]3d104s24p2
36. Which of the following ions have the same ground state electron configuration: S2–, N3–, Mg2+, and Br–? A. N3– and Mg2+ B. S2–, N3–, and Br– C. S2– and Br– D. Mg2+ and Br– E. S2–, N3–, Mg2+, and Br–
37. Hund's rule predicts that A. the most stable electronic structure of an atom has electron spins paired. B. no two electrons will share the same orbital. C. electrons in an orbital have equal but opposite charges. D. electrons must have opposite spins to share an orbital. E. when several orbitals of equal energy are available, as in a given subshell, electrons enter singly with parallel spins.
38. What is the symbol of the atom or ion with the following orbital diagram?
1s (¯)
A. 5B B. 7N C. 8O D. 16S E. 9F
2s (¯)
2p ( )( )( )
39. What is the symbol of the atom or ion with the following orbital diagram?
1s (¯)
2s (¯)
2p (¯)(¯)(¯)
3s (¯)
3p (¯)(¯)(¯)
A. 26Fe B. 26Fe2+ C. 27Co+ D. 28Ni E. 28Ni2+ 40. What is the symbol of an ion with the following orbital diagram?
3d (¯)( )( )( )( )
[Ar]
4s ( )
A. 25Mn2+ B. 28Ni2+ C. 26Fe3+ D. 24Cr3+ E. 27Co3+ 41. What is the correct orbital diagram for phosphorus?
1) 2) 3) 4) 5)
A. 1 B. 2 C. 3 D. 4 E. 5
1s (¯) (¯) (¯) (¯) (¯)
2s (¯) (¯) (¯) (¯) (¯)
2p ( )( )( ) (¯)( )( ) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯)
3s
3p
(¯) (¯) ( )
(¯)( )( ) ( )( )( ) (¯)( )( )
3d (¯)(¯)(¯)( )( )
4s ( )
42. What is the correct orbital diagram for Ca2+?
1) 2) 3) 4) 5)
1s (¯) (¯) (¯) (¯) (¯)
2s (¯) (¯) (¯) (¯) (¯)
2p (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯)
3s
3p
4s
(¯) (¯) (¯) (¯)
(¯)(¯)( ) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯)
( ) (¯)
A. 1 B. 2 C. 3 D. 4 E. 5 43. What is the correct orbital diagram for Fe2+?
1) 2) 3) 4) 5)
[Ar] [Ar] [Ar] [Ar] [Ar]
3d ( )( )( )( )( ) ( )( )( )( )( ) ( )( )( )( )( ) (¯)( )( )( )( ) (¯)(¯)( )( )( )
A. 1 B. 2 C. 3 D. 4 E. 5 44. What is the correct orbital diagram for N3-?
1) 2) 3) 4) 5)
A. 1 B. 2 C. 3 D. 4 E. 5
1s (¯) (¯) (¯) (¯) (¯)
2s (¯) (¯) (¯) (¯) ( )
2p ( )( )( ) ( )( )( ) ( )( )( ) (¯)(¯)(¯) (¯)(¯)(¯)
4s (¯) ( ) ( ) ( ) ( )
45. Elements and compounds with unpaired electrons are attracted to a magnetic field. These materials are called paramagnetic. Which of the following ions are paramagnetic in the ground state? A. IB. Ni2+ C. P3D. Ca2+ E. Ti4+
46. Rank F, Cl, and Br in order of increasing first ionization energy. A. F < Cl < Br B. Cl < F < Br C. Cl < Br < F D. Br < F < Cl E. Br < Cl < F
47. In general, atomic radii A. increase down a group and decrease across a period. B. increase down a group and increase across a period. C. decrease down a group and decrease across a period. D. are proportional to atomic mass. E. decrease down a group and increase across a period.
48. Place the following atoms in order of increasing atomic radius: Al, Cl, Mg, O, and P. A. Cl < O < P < Al < Mg B. Cl < P < Al < Mg < O C. O < Cl < P < Al < Mg D. O < Mg < Al < P < Cl E. none of the above
49. Place the following atoms in order of increasing atomic radii: K, Na, Be, and Li? A. Li < Na < K < Be B. Be < K < Na < Li C. Be < Li < Na < K D. K < Na < Li < Be E. Li < Be < Na < K
50. Place the following ions in order of increasing radius: Al3+, F-, Mg2+, and N3-. A. F- < Mg2+ < N3- < Al3+ B. F- < N3- < Al3+ < Mg2+ C. F- < N3- < Mg2+ < Al3+ D. N3- < F- < Mg2+ < Al3+ E. Al3+ < Mg2+ < F- < N3-
51. Place the following atoms or ions in order of increasing radius: O, O2-, S2-, and Se2-. A. O2- < O < S2- < Se2B. Se2- < S2- < O2- < O C. O2- < O < Se2- < S2D. Se2- < S2- < O < O2E. O < O2- < S2- < Se2-
52. Which of the following chemical expressions refers to the first ionization energy of calcium? A. Ca(s) ® Ca+(s) + eB. Ca(g) + e- ® Ca+(g) C. Ca(s) + e- ® Ca+(s) D. Ca(g) ® Ca+(g) + eE. Ca(g) ® Ca2+(g) + e-
53. Which of the following elements is assigned an electronegativity value of 4.0, which is the greatest electronegativity value? A. H B. Au C. F D. He E. Cs
54. Place the following atoms in order of increasing ionization energy: N, O, and P. A. N < P < O B. O < P < N C. P < N < O D. N < O < P E. P < O < N
55. Electronegativity increases A. moving down a group in the periodic table. B. moving from left to right across the periodic table. C. with increasing atomic mass. D. when electrons are paired. E. with increasing atomic radii.
56. The color of the shell of a cooked lobster is due to A. the denaturing of the protein crustacyanin, releasing astaxanthin a pigment that absorbs blue light. B. the denaturing of the protein crustacyanin, releasing astaxanthin a pigment that absorbs red light. C. the temperature of the shell. D. the reaction of the water with the pigments in the lobster shell. E. the pigments dissolved in the water absorbing into the lobster shell.
57. Which of the five atoms Na, N, Cl, Mg, or Al has the largest atomic radius? A. Na B. N C. Cl D. Mg E. Al
58. Place the following ions in order of increasing size: Al3+, P3-, and S2A. P3- < S2- < Al3+ B. Al3+ < P3-< S2C. Al3+ < S2- < P3D. P3- < Al3+ < S2E. S2- < Al3+ < P3-
59. Glen Seaborg is known for the discovery of many elements. What do these elements have in common? A. they all occur in nature B. they all contain 5f electrons C. they are all isotopes of uranium D. they are all required to prepare a nuclear bomb E. they are all used in the electronics industry
Chapter 6--Electronic Structure and the Periodic Table Key
1. The Navy uses electromagnetic radiation of extremely long wavelengths to communicate with submerged submarines. The Navy's ELF (Extremely Low Frequency) systems emit 76 s-1 radiation from facilities located in Wisconsin and Michigan. What is the wavelength of this radiation (in meters)? A. 2.1 ´ 10-7 m B. 1.3 ´ 10-3 m C. 2.5 ´ 103 m D. 3.9 ´ 106 m E. 2.3 ´ 1010 m
2. If a radio wave has a frequency of 17.0 MHz, what is the wavelength of this radiation? A. 5.10 ´ 1015 m B. 5.67 ´ 10–2 m C. 5.88 ´ 10–2 m D. 1.76 ´ 101 m E. 1.76 ´ 106 m
3. If a cordless phone operates at a frequency of 9.00 ´ 108 s–1. What is the wavelength of this radiation (in nm)? A. 0.333 m B. 5.96 ´ 10–25 m C. 3.33 ´ 10–9 m D. 3.71 ´ 10–18 m E. 2.70 ´ 1017 m
4. What is the frequency of a gamma ray radiation that has a wavelength of 11.4 pm? A. 3.80 ´ 10–20 s–1 B. 1.74 ´ 10–14 s–1 C. 2.63 ´ 107 s–1 D. 3.42 ´ 109 s–1 E. 2.63 ´ 1019 s–1
5. Green laser pointers emit radiation at 532 nm. What is the frequency of this radiation? A. 8.12 ´ 1013 Hz B. 5.64 ´ 1014 Hz C. 1.60 ´ 1015 Hz D. 9.10 ´ 1015 Hz E. 1.60 ´ 1016 Hz
6. Which of the following regions of the electromagnetic spectrum have longer wavelengths than visible light? 1. 2. 3.
infrared radiation ultraviolet radiation microwave radiation
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1 and 3 7. Place the following regions of the electromagnetic spectrum in order from highest to lowest energy. A. infrared > gamma rays > x-rays > visible > radio B. gamma rays > x-rays > visible > infrared > radio C. x-rays > gamma rays > infrared > visible > radio D. radio > x-rays > gamma rays > visible > infrared E. visible > infrared > radio > x-rays > gamma rays
8. The ____ of a photon of light is ____ proportional to its frequency and ____ proportional to its wavelength. A. energy, directly, inversely B. energy, inversely, directly C. velocity, directly, inversely D. intensity, inversely, directly E. amplitude, directly, inversely
9. Some digital cordless phones operate at 2.4 GHz (1 GHz = 109 Hz). What is the energy, in joules, of a single 2.4 GHz photon? A. 8.3 ´ 10-35 J B. 1.6 ´ 10-24 J C. 5.4 ´ 10-19 J D. 0.96 J E. 0.13 J
10. Excited sodium atoms emit light in the infrared at 589 nm. What is the energy of a single photon with this wavelength? A. 5.09 ´ 1014 J B. 1.12 ´ 10–27 J C. 3.37 ´ 10–19 J D. 3.37 ´ 10–28 J E. 1.30 ´ 10–-19 J
11. When Cs-137 decays, it emits gamma radiation. The energy of one photon is 1.06 ´ 10-13 J. What is the wavelength of this radiation? A. 1.87 ´ 10-12 m B. 2.08 ´ 10-12 m C. 2.44 ´ 10-12 m D. 1.88 ´ 10-11 m E. 1.58 ´ 10-10 m
12. If the energy of 1.00 mole of photons is 245 kJ, what is the wavelength of the light? A. 122 nm B. 488 nm C. 1220 nm D. 787 nm E. 811 nm
13. According to the Bohr model for the hydrogen atom, the energy necessary to excite an electron from n = 2 to n = 3 is ____ the energy necessary to excite an electron from n = 3 to n = 4. A. greater than B. equal to C. less than D. either equal to or greater than E. either less than or equal to
14. Calculate the wavelength of the line in the Lyman series that results from the transition n = 3 to n = 1. The Rydberg constant equals 2.180 ´ 10-18 J. A. 45.59 nm B. 72.81 nm C. 91.12 nm D. 102.5 nm E. 136.7 nm
15. Calculate the energy of a photon in the Balmer series that results from the transition n = 3 to n = 2. What is the region of the electromagnetic spectrum? The Rydberg constant equals 2.180 ´ 10-18 J. A. 3.028 ´ 10-19 J; visible B. 4.578 ´ 10-19 J; infrared C. 3.633 ´ 10-19 J; visible D. 3.633 ´ 10-18 J; ultraviolet E. 2.180 ´ 10-18 J; ultraviolet
16. A line in the Pfund series (nlo = 5) occurs at 3.74 ´ 10-6 m. What is nhi for this transition? The Rydberg constant equals 2.180 ´ 10-18 J. A. 6 B. 7 C. 8 D. 9 E. 10
17. For which of the following transitions would a hydrogen atom absorb a photon with the longest wavelength? A. n = 1 to n = 2 B. n = 4 to n = 6 C. n = 5 to n = 4 D. n = 7 to n = 6 E. n = 6 to n = 7
18. The Schrödinger wave equation A. proves electrons have positive and negative spins. B. calculates the precise position and momentum of an electron at any given time. C. is used to compute the wavelength of small particles. D. can be solved to find the probability of finding an electron in a region of space. E. proves that photons are particles.
19. The Pauli exclusion principle states that A. no two electrons from a given atom can have the same spin. B. no two electrons from a given atom can have the same four quantum numbers. C. two electrons can occupy an orbital if they have the same spin. D. two electrons can occupy an orbital if they have opposite spins. E. two electrons can occupy an orbital if they have opposite charges.
20. What type of orbital is designated n = 3, A. 3s B. 3p C. 3d D. 2f E. 2d
= 2,
= –2?
21. Which of the following sets of quantum numbers refers to a 2s orbital? A. n = 1,
= 2,
= 2, ms = +
B. n = 1,
= 2,
= 1, ms = +
C. n = 2,
= 2,
= 0, ms = +
D. n = 2,
= 1,
= -1, ms = +
E. n = 2,
= 0,
= 0, ms = +
22. All of the following sets of quantum numbers are allowed EXCEPT A. n = 1,
= 0,
= 1, ms = -
.
B. n = 2,
= 1,
= 0, ms = +
.
C. n = 3,
= 1,
= -1, ms = -
.
D. n = 4,
= 3,
= -1, ms = -
.
E. n = 6,
= 3,
= -3, ms = +
.
23. What is the total number of orbitals having n = 4 and A. 3 B. 5 C. 7 D. 9 E. 10
= 2?
24. How many orbitals have the following set of quantum numbers: n = 5, A. 0 B. 1 C. 3 D. 6 E. 7
25. What is the total capacity of electrons in n = 3? A. 5 B. 10 C. 18 D. 32 E. 50
26. What is the total capacity of electrons in n = 5, A. 2 B. 6 C. 10 D. 14 E. 32
= 3?
27. Which of the following sets of quantum numbers refers to a 4f orbital? A. n = 4,
= 1,
= 0, ms = +
B. n = 4,
= 2,
= -1, ms = -
C. n = 4,
= 3,
= -2, ms = -
D. n = 4,
= 4,
= 0, ms = +
E. n = 4,
= 4,
= -4, ms = +
28. What type of orbital is designated by n = 5, A. 5d B. 2f C. 5p D. 2s E. 5p
= 2,
= +1?
= 3,
= +2?
29. Which of the following properties is associated with the value of the A. the shape of an orbital B. the size of an orbital C. the number of electrons in an orbital D. the energy of an orbital E. the orientation in space of an orbital
30. What is the ground state electron configuration of 25Mn? A. 1s22s22p63s23p63d64s1 B. 1s22s22p63s23p63d54s2 C. 1s22s22p63s23p63d7 D. 1s22s22p63s23p63d5 E. 1s22s22p63s23p64s24d5
31. What is the electron configuration of 80Hg? A. [Xe]6s26d10 B. [Kr]5d106s2 C. [Xe]4f145d106s2 D. [Xe]5d106s2 E. [Kr]4d104f146s2
32. What is the ground state electron configuration of 94Pu? A. [Xe]5f66s2 B. [Xe]6f67s2 C. [Rn]6f67s2 D. [Rn]5f67s2 E. [Rn]5f66d107s2
33. What is the electron configuration of S2-? A. 1s22s22p6 B. 1s22s22p63s2 C. 1s22s22p63s23p4 D. 1s22s22p63s23p6 E. 1s22s22p63s23p64s1
quantum number?
34. What is the ground state electron configuration for Cr3+? A. [Ar] B. [Ar]3d74s2 C. [Ar]3d14s2 D. [Ar]3d24s1 E. [Ar]3d3
35. What is the electron configuration of Ga+? A. [Ar]3d104s2 B. [Ar]3d104p2 C. [Ar]3d84s24p2 D. [Ar]3d104s24d4 E. [Ar]3d104s24p2
36. Which of the following ions have the same ground state electron configuration: S2–, N3–, Mg2+, and Br–? A. N3– and Mg2+ B. S2–, N3–, and Br– C. S2– and Br– D. Mg2+ and Br– E. S2–, N3–, Mg2+, and Br–
37. Hund's rule predicts that A. the most stable electronic structure of an atom has electron spins paired. B. no two electrons will share the same orbital. C. electrons in an orbital have equal but opposite charges. D. electrons must have opposite spins to share an orbital. E. when several orbitals of equal energy are available, as in a given subshell, electrons enter singly with parallel spins.
38. What is the symbol of the atom or ion with the following orbital diagram?
1s (¯)
A. 5B B. 7N C. 8O D. 16S E. 9F
2s (¯)
2p ( )( )( )
39. What is the symbol of the atom or ion with the following orbital diagram?
1s (¯)
2s (¯)
2p (¯)(¯)(¯)
3s (¯)
3p (¯)(¯)(¯)
A. 26Fe B. 26Fe2+ C. 27Co+ D. 28Ni E. 28Ni2+ 40. What is the symbol of an ion with the following orbital diagram?
3d (¯)( )( )( )( )
[Ar]
4s ( )
A. 25Mn2+ B. 28Ni2+ C. 26Fe3+ D. 24Cr3+ E. 27Co3+ 41. What is the correct orbital diagram for phosphorus?
1) 2) 3) 4) 5)
A. 1 B. 2 C. 3 D. 4 E. 5
1s (¯) (¯) (¯) (¯) (¯)
2s (¯) (¯) (¯) (¯) (¯)
2p ( )( )( ) (¯)( )( ) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯)
3s
3p
(¯) (¯) ( )
(¯)( )( ) ( )( )( ) (¯)( )( )
3d (¯)(¯)(¯)( )( )
4s ( )
42. What is the correct orbital diagram for Ca2+?
1) 2) 3) 4) 5)
1s (¯) (¯) (¯) (¯) (¯)
2s (¯) (¯) (¯) (¯) (¯)
2p (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯)
3s
3p
4s
(¯) (¯) (¯) (¯)
(¯)(¯)( ) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯)
( ) (¯)
A. 1 B. 2 C. 3 D. 4 E. 5 43. What is the correct orbital diagram for Fe2+?
1) 2) 3) 4) 5)
[Ar] [Ar] [Ar] [Ar] [Ar]
3d ( )( )( )( )( ) ( )( )( )( )( ) ( )( )( )( )( ) (¯)( )( )( )( ) (¯)(¯)( )( )( )
A. 1 B. 2 C. 3 D. 4 E. 5 44. What is the correct orbital diagram for N3-?
1) 2) 3) 4) 5)
A. 1 B. 2 C. 3 D. 4 E. 5
1s (¯) (¯) (¯) (¯) (¯)
2s (¯) (¯) (¯) (¯) ( )
2p ( )( )( ) ( )( )( ) ( )( )( ) (¯)(¯)(¯) (¯)(¯)(¯)
4s (¯) ( ) ( ) ( ) ( )
45. Elements and compounds with unpaired electrons are attracted to a magnetic field. These materials are called paramagnetic. Which of the following ions are paramagnetic in the ground state? A. IB. Ni2+ C. P3D. Ca2+ E. Ti4+
46. Rank F, Cl, and Br in order of increasing first ionization energy. A. F < Cl < Br B. Cl < F < Br C. Cl < Br < F D. Br < F < Cl E. Br < Cl < F
47. In general, atomic radii A. increase down a group and decrease across a period. B. increase down a group and increase across a period. C. decrease down a group and decrease across a period. D. are proportional to atomic mass. E. decrease down a group and increase across a period.
48. Place the following atoms in order of increasing atomic radius: Al, Cl, Mg, O, and P. A. Cl < O < P < Al < Mg B. Cl < P < Al < Mg < O C. O < Cl < P < Al < Mg D. O < Mg < Al < P < Cl E. none of the above
49. Place the following atoms in order of increasing atomic radii: K, Na, Be, and Li? A. Li < Na < K < Be B. Be < K < Na < Li C. Be < Li < Na < K D. K < Na < Li < Be E. Li < Be < Na < K
50. Place the following ions in order of increasing radius: Al3+, F-, Mg2+, and N3-. A. F- < Mg2+ < N3- < Al3+ B. F- < N3- < Al3+ < Mg2+ C. F- < N3- < Mg2+ < Al3+ D. N3- < F- < Mg2+ < Al3+ E. Al3+ < Mg2+ < F- < N3-
51. Place the following atoms or ions in order of increasing radius: O, O2-, S2-, and Se2-. A. O2- < O < S2- < Se2B. Se2- < S2- < O2- < O C. O2- < O < Se2- < S2D. Se2- < S2- < O < O2E. O < O2- < S2- < Se2-
52. Which of the following chemical expressions refers to the first ionization energy of calcium? A. Ca(s) ® Ca+(s) + eB. Ca(g) + e- ® Ca+(g) C. Ca(s) + e- ® Ca+(s) D. Ca(g) ® Ca+(g) + eE. Ca(g) ® Ca2+(g) + e-
53. Which of the following elements is assigned an electronegativity value of 4.0, which is the greatest electronegativity value? A. H B. Au C. F D. He E. Cs
54. Place the following atoms in order of increasing ionization energy: N, O, and P. A. N < P < O B. O < P < N C. P < N < O D. N < O < P E. P < O < N
55. Electronegativity increases A. moving down a group in the periodic table. B. moving from left to right across the periodic table. C. with increasing atomic mass. D. when electrons are paired. E. with increasing atomic radii.
56. The color of the shell of a cooked lobster is due to A. the denaturing of the protein crustacyanin, releasing astaxanthin a pigment that absorbs blue light. B. the denaturing of the protein crustacyanin, releasing astaxanthin a pigment that absorbs red light. C. the temperature of the shell. D. the reaction of the water with the pigments in the lobster shell. E. the pigments dissolved in the water absorbing into the lobster shell.
57. Which of the five atoms Na, N, Cl, Mg, or Al has the largest atomic radius? A. Na B. N C. Cl D. Mg E. Al
58. Place the following ions in order of increasing size: Al3+, P3-, and S2A. P3- < S2- < Al3+ B. Al3+ < P3-< S2C. Al3+ < S2- < P3D. P3- < Al3+ < S2E. S2- < Al3+ < P3-
59. Glen Seaborg is known for the discovery of many elements. What do these elements have in common? A. they all occur in nature B. they all contain 5f electrons C. they are all isotopes of uranium D. they are all required to prepare a nuclear bomb E. they are all used in the electronics industry
Chapter 7--Covalent Bonding 1. A pair of electrons that is shared between two atoms is A. a covalent bond. B. a lone pair. C. a double bond. D. an ionic bond. E. both a covalent bond and a double bond.
2. What is the expected number of valence electrons for an element in group 16? A. 0 B. 2 C. 4 D. 6 E. 16
3. A nitrogen atom has _____ valence electrons. A. 0 B. 3 C. 5 D. 7 E. 8
4. A fluoride ion, F–, has _____ valence electrons. A. 6 B. 7 C. 8 D. 9 E. 10
5. Which of the following is a correct Lewis structure for oxygen? A. B. C. D. E.
6. Which of the following molecules or ions will have a Lewis structure most like that of phosphorus trichloride, PCl3? A. ClO3– B. SO3 C. CO32– D. BF3 E. Cl2CO
7. Which of the following is/are possible Lewis structures for C2H6O?
A. 1 B. 2 C. 3 D. 2 and 3 E. 1, 2, and 3
8. All of the following Lewis structures of nitrogen oxides are possible EXCEPT
A. N2O. B. N2O4. C. N2O3. D. N2O5. E. All of the above are correct structures.
9. Which of the following is a correct Lewis structure for sulfur dioxide, SO2? A. B. C. D. E.
10. Which of the following is a correct Lewis structure for PH3?
A.
B.
C.
D.
E.
11. Which of the following elements is able to form a molecular structure that exceeds the octet rule? A. C B. B C. N D. F E. S
12. What is the correct Lewis structure of SF4?
A.
B.
C.
D.
E.
13. What is the correct Lewis structure for IF3?
A.
B.
C.
D.
E.
14. How many different molecules have the formula C5H12? A. 1 B. 2 C. 3 D. 4 E. 6
15. Which of the following are correct resonance structures of SO3?
A. (1) and (5) B. (2) and (4) C. (1), (2), and (4) D. (2), (3) and (4) E. (1), (2), (4), and (5)
16. How many possible resonance structures exist for carbonate ion, CO32–? A. 0 B. 1 C. 3 D. 6 E. 8
17. How many possible resonance structures exist for the formate ion, HCO2-? A. 0 B. 2 C. 3 D. 4 E. 8
18. Which of the following are correct resonance structures of N2O4?
A. (1) and (2) B. (2) and (3) C. (1), (2), and (3) D. (2), (3), and (4) E. (1), (2), (3), and (4)
19. Which of the following species will have a Lewis structure with a molecular geometry similar to IF4–? A. XeF4 B. SO42– C. PF4+ D. SF4 E. IO4–
20. Which of the following species has a Lewis structure with a molecular geometry similar to SO2? A. H2S B. NO2C. NO2 D. N2O E. ClO2-
21. Which of the following species has a Lewis structure with a molecular geometry similar to SF4? A. BrF4+ B. ICl4C. NH4+ D. SO42E. CCl4
22. Which of the following species has a Lewis structure with a molecular geometry similar to SO3? A. NH3 B. ICl3 C. CO32D. SO32E. PCl3
23. How many lone pairs of electrons are on the sulfur atom in sulfite ion, SO32-? A. 0 B. 1 C. 2 D. 3 E. 4
24. The central atom in SCl2 is surrounded by A. two single bonds and no lone pairs of electrons. B. two single bonds and one lone pair of electrons. C. two single bonds and two lone pairs of electrons. D. one single bond, one double bond, and no lone pairs of electrons. E. one single bond, one double bond, and one lone pair of electrons.
25. The nitrogen atom in cyanide ion, CN-, is surrounded by A. one single bond and three lone pairs of electrons. B. one double bond and one lone pair of electrons. C. one double bond and two lone pairs of electrons. D. one triple bond and one lone pair of electrons. E. one triple bond and no lone pairs of electrons.
26. Formal charge is A. the absolute value of the charge on a polyatomic anion or cation. B. the difference between the number of lone pairs of electrons and shared pairs of electrons on any atom in a Lewis structure. C. the difference between the number of valence electrons and the number of protons in any given atom. D. equal to the number of valence electrons in a free atom minus the number of shared in covalent bonds. E. the difference between the number of valence electrons in a free atom and the number of electrons assigned to the atom in a Lewis structure.
27. A Lewis structure of OCl– ion is drawn below. What is the formal charge on each atom?
A. Cl atom = +1 and each O atom = –1 B. Cl atom = 0 and each O atom = –1 C. Cl atom = –1 and each O atom = 0 D. Cl atom = +3 and each O atom = –2 E. Cl atom = 0, one O atom = 0, one O atom = –1
28. What is the formal charge on each atom in CN-? A. C = 0, N = 0 B. C = +1, N = -1 C. C = -1, N = 0 D. C = +2, N = -3 E. C = +4, N = -5
29. Using formal charges and the octet rule, determine which Lewis structure of OCN- is most stable.
A.
B.
C.
D. E.
30. Use VSEPR theory to predict the molecular geometry of nitrogen trichloride, NCl3. A. linear B. trigonal planar C. bent D. tetrahedron E. trigonal pyramid
31. Use VSEPR theory to predict the molecular geometry of ClO3-. A. bent B. tetrahedron C. square planar D. trigonal planar E. trigonal pyramid
32. Use VSEPR theory to predict the molecular geometry of ICl3. A. trigonal planar B. trigonal pyramid C. trigonal bipyramid D. t-shaped E. octahedron
33. Use VSEPR theory to predict the molecular geometry of SO2. A. bent B. linear C. trigonal planar D. tetrahedron E. trigonal pyramid
34. Use VSEPR theory to predict the molecular geometry of phosgene, Cl2CO. A. t-shaped B. trigonal planar C. trigonal bipyramid D. square planar E. tetrahedronl
35. Use VSEPR theory to predict the molecular geometry of IF5. A. octahedron B. square planar C. tetrahedron D. see-saw E. square pyramid
36. Use VSEPR theory to predict the molecular geometry of I3–. A. bent B. linear C. trigonal planar D. t-shaped E. octahedron
37. Use VSEPR theory to predict the molecular geometry of H2Se. A. bent B. linear C. tetrahedron D. trigonal planar E. trigonal pyramid
38. Use VSEPR theory to predict the molecular geometry of NH4+. A. trigonal pyramid B. square pyramid C. see-saw D. tetrahedron E. trigonal planar
39. Use VSEPR theory to predict the molecular geometry around each carbon atom in acetylene, C 2H2. A. linear B. bent C. trigonal planar D. tetrahedron E. octahedron
40. Which of the following species have the same molecular geometry: H3O+, H2CO, NH3, and ICl3? A. H3O+ and H2CO B. H2CO and ICl3 C. H3O+ and NH3 D. H3O+, NH3, and ICl3 E. None of the species have the same molecular geometry.
41. Which of the following species have the same molecular geometry: PO43-, SF4, PF5, and XeF4? A. PO43- and SF4 B. SF4 and PF5 C. PO43-, SF4, and XeF4 D. SF4 and XeF4 E. None of the species have the same molecular geometry.
42. Which of the following species have the same molecular geometry: CO2, H2O, BeCl2, and N2O? A. CO2 and N2O only B. H2O and N2O only C. H2O and BeCl2 only D. CO2 and BeCl2 only E. CO2, BeCl2, and N2O
43. Which of the following species have the same molecular geometry: XeF4, ClF4+, SF4, PO43-? A. XeF4 and SF4 B. ClF4+, and SF4 C. ClF4+ and PO43D. XeF4 and ClF4+ E. XeF4, SF4, and PO43-
44. What are the O–Cl–O bond angles in ClO4-? A. 90 B. 109.5 C. 120 D. 90 and 120 E. 180
45. What are the approximate O–N–O bond angles in a nitrate ion? A. 90 B. 109.5 C. 120 D. 90 and 120 E. 180
46. What are the F–Xe–F bond angles in XeF4? A. 90 and 180 B. 109.5 C. 120 D. 60 and 120 E. 180
47. What are the approximate bond angles in SF4? A. 90 B. 109.5 C. 120 D. 90 and 120 E. 180
48. Which of the following molecules has the smallest bond angle between any two hydrogen atoms? A. CH4 B. H2O C. BH3 D. PH3 E. SiH4
49. Which of the following compounds has polar covalent bonds: CCl4, Cl2, HCl, and KCl? A. CCl4 only B. Cl2 only C. HCl and KCl D. Cl2 and KCl E. CCl4 and HCl
50. Which of the following molecules are polar: H2S, CO2, NH3, BH3, and CCl4? A. BH3 B. H2S and NH3 C. H2S, CO2, and CCl4 D. CO2, NH3, and CCl4 E. NH3, BH3 and CCl4
51. What is the hybridization of the carbon atoms in benzene, C6H6? A. sp B. sp2 C. sp3 D. sp3d E. sp3d2
52. What is the hybridization of the phosphorus atom in PCl3? A. sp B. sp2 C. sp3 D. sp3d E. sp3d2
53. What is the hybridization of the central atom in I3-? A. sp B. sp2 C. sp3 D. sp3d E. sp3d2
54. What is the hybridization of the central atom in SO2? A. sp B. sp2 C. sp3 D. sp3d E. sp3d2
55. What hybridization change does the carbon atom undergo in the combustion of methane? CH4(g) + 2O2(g) ® CO2(g) + 2H2O(g) A. sp ® sp2 B. sp2 ® sp3 C. sp3 ® sp D. sp2 ® sp E. none
56. How many sigma and pi bonds are present in the following molecule?
A. 8 sigma bonds and 1 pi bond B. 8 sigma bonds and 2 pi bonds C. 10 sigma bonds and 2 pi bonds D. 11 sigma bonds and 2 pi bonds E. 11 sigma bonds and 1 pi bond
57. How many sigma and pi bonds are present in H2CO? A. 1 sigma bond and 3 pi bonds B. 2 sigma bonds and 2 pi bonds C. 2 sigma bonds and 1 pi bond D. 3 sigma bonds and 1 pi bond E. none of the above
58. G. N. Lewis is well known for multiple contributions to the study of chemistry. Which of these below is attributed to Lewis? A. covalent bonding B. ionic bonding C. VSEPR D. hybridization E. paramagnetism
59. This gas changed the perception of the reactivity of this group. A. oxygen B. xenon C. helium D. hydrogen E. chlorine
60. Write the singly bonded Lewis dot structure for BF3. Which of the following statements best describes this structure? A. It obeys the octet rule on all atoms. B. It has less than an octet on at least one atom. C. It has a lone pair of electrons on the boron atom. D. It has less than an octet of electrons on all atoms. E. It exceeds the octet rule.
61. Write the singly bonded Lewis dot structure for SF6. Which of the following statements best describes this structure? A. It obeys the octet rule on all atoms. B. It has less than an octet on at least one atom. C. It has a lone pair of electrons on the sulfur atom. D. It has less than an octet of electrons on all atoms. E. It exceeds the octet rule.
Chapter 7--Covalent Bonding Key
1. A pair of electrons that is shared between two atoms is A. a covalent bond. B. a lone pair. C. a double bond. D. an ionic bond. E. both a covalent bond and a double bond.
2. What is the expected number of valence electrons for an element in group 16? A. 0 B. 2 C. 4 D. 6 E. 16
3. A nitrogen atom has _____ valence electrons. A. 0 B. 3 C. 5 D. 7 E. 8
4. A fluoride ion, F–, has _____ valence electrons. A. 6 B. 7 C. 8 D. 9 E. 10
5. Which of the following is a correct Lewis structure for oxygen? A. B. C. D. E.
6. Which of the following molecules or ions will have a Lewis structure most like that of phosphorus trichloride, PCl3? A. ClO3– B. SO3 C. CO32– D. BF3 E. Cl2CO
7. Which of the following is/are possible Lewis structures for C2H6O?
A. 1 B. 2 C. 3 D. 2 and 3 E. 1, 2, and 3
8. All of the following Lewis structures of nitrogen oxides are possible EXCEPT
A. N2O. B. N2O4. C. N2O3. D. N2O5. E. All of the above are correct structures.
9. Which of the following is a correct Lewis structure for sulfur dioxide, SO2? A. B. C. D. E.
10. Which of the following is a correct Lewis structure for PH3?
A.
B.
C.
D.
E.
11. Which of the following elements is able to form a molecular structure that exceeds the octet rule? A. C B. B C. N D. F E. S
12. What is the correct Lewis structure of SF4?
A.
B.
C.
D.
E.
13. What is the correct Lewis structure for IF3?
A.
B.
C.
D.
E.
14. How many different molecules have the formula C5H12? A. 1 B. 2 C. 3 D. 4 E. 6
15. Which of the following are correct resonance structures of SO3?
A. (1) and (5) B. (2) and (4) C. (1), (2), and (4) D. (2), (3) and (4) E. (1), (2), (4), and (5)
16. How many possible resonance structures exist for carbonate ion, CO32–? A. 0 B. 1 C. 3 D. 6 E. 8
17. How many possible resonance structures exist for the formate ion, HCO2-? A. 0 B. 2 C. 3 D. 4 E. 8
18. Which of the following are correct resonance structures of N2O4?
A. (1) and (2) B. (2) and (3) C. (1), (2), and (3) D. (2), (3), and (4) E. (1), (2), (3), and (4)
19. Which of the following species will have a Lewis structure with a molecular geometry similar to IF4–? A. XeF4 B. SO42– C. PF4+ D. SF4 E. IO4–
20. Which of the following species has a Lewis structure with a molecular geometry similar to SO2? A. H2S B. NO2C. NO2 D. N2O E. ClO2-
21. Which of the following species has a Lewis structure with a molecular geometry similar to SF4? A. BrF4+ B. ICl4C. NH4+ D. SO42E. CCl4
22. Which of the following species has a Lewis structure with a molecular geometry similar to SO3? A. NH3 B. ICl3 C. CO32D. SO32E. PCl3
23. How many lone pairs of electrons are on the sulfur atom in sulfite ion, SO32-? A. 0 B. 1 C. 2 D. 3 E. 4
24. The central atom in SCl2 is surrounded by A. two single bonds and no lone pairs of electrons. B. two single bonds and one lone pair of electrons. C. two single bonds and two lone pairs of electrons. D. one single bond, one double bond, and no lone pairs of electrons. E. one single bond, one double bond, and one lone pair of electrons.
25. The nitrogen atom in cyanide ion, CN-, is surrounded by A. one single bond and three lone pairs of electrons. B. one double bond and one lone pair of electrons. C. one double bond and two lone pairs of electrons. D. one triple bond and one lone pair of electrons. E. one triple bond and no lone pairs of electrons.
26. Formal charge is A. the absolute value of the charge on a polyatomic anion or cation. B. the difference between the number of lone pairs of electrons and shared pairs of electrons on any atom in a Lewis structure. C. the difference between the number of valence electrons and the number of protons in any given atom. D. equal to the number of valence electrons in a free atom minus the number of shared in covalent bonds. E. the difference between the number of valence electrons in a free atom and the number of electrons assigned to the atom in a Lewis structure.
27. A Lewis structure of OCl– ion is drawn below. What is the formal charge on each atom?
A. Cl atom = +1 and each O atom = –1 B. Cl atom = 0 and each O atom = –1 C. Cl atom = –1 and each O atom = 0 D. Cl atom = +3 and each O atom = –2 E. Cl atom = 0, one O atom = 0, one O atom = –1
28. What is the formal charge on each atom in CN-? A. C = 0, N = 0 B. C = +1, N = -1 C. C = -1, N = 0 D. C = +2, N = -3 E. C = +4, N = -5
29. Using formal charges and the octet rule, determine which Lewis structure of OCN- is most stable.
A.
B.
C.
D. E.
30. Use VSEPR theory to predict the molecular geometry of nitrogen trichloride, NCl3. A. linear B. trigonal planar C. bent D. tetrahedron E. trigonal pyramid
31. Use VSEPR theory to predict the molecular geometry of ClO3-. A. bent B. tetrahedron C. square planar D. trigonal planar E. trigonal pyramid
32. Use VSEPR theory to predict the molecular geometry of ICl3. A. trigonal planar B. trigonal pyramid C. trigonal bipyramid D. t-shaped E. octahedron
33. Use VSEPR theory to predict the molecular geometry of SO2. A. bent B. linear C. trigonal planar D. tetrahedron E. trigonal pyramid
34. Use VSEPR theory to predict the molecular geometry of phosgene, Cl2CO. A. t-shaped B. trigonal planar C. trigonal bipyramid D. square planar E. tetrahedronl
35. Use VSEPR theory to predict the molecular geometry of IF5. A. octahedron B. square planar C. tetrahedron D. see-saw E. square pyramid
36. Use VSEPR theory to predict the molecular geometry of I3–. A. bent B. linear C. trigonal planar D. t-shaped E. octahedron
37. Use VSEPR theory to predict the molecular geometry of H2Se. A. bent B. linear C. tetrahedron D. trigonal planar E. trigonal pyramid
38. Use VSEPR theory to predict the molecular geometry of NH4+. A. trigonal pyramid B. square pyramid C. see-saw D. tetrahedron E. trigonal planar
39. Use VSEPR theory to predict the molecular geometry around each carbon atom in acetylene, C 2H2. A. linear B. bent C. trigonal planar D. tetrahedron E. octahedron
40. Which of the following species have the same molecular geometry: H3O+, H2CO, NH3, and ICl3? A. H3O+ and H2CO B. H2CO and ICl3 C. H3O+ and NH3 D. H3O+, NH3, and ICl3 E. None of the species have the same molecular geometry.
41. Which of the following species have the same molecular geometry: PO43-, SF4, PF5, and XeF4? A. PO43- and SF4 B. SF4 and PF5 C. PO43-, SF4, and XeF4 D. SF4 and XeF4 E. None of the species have the same molecular geometry.
42. Which of the following species have the same molecular geometry: CO2, H2O, BeCl2, and N2O? A. CO2 and N2O only B. H2O and N2O only C. H2O and BeCl2 only D. CO2 and BeCl2 only E. CO2, BeCl2, and N2O
43. Which of the following species have the same molecular geometry: XeF4, ClF4+, SF4, PO43-? A. XeF4 and SF4 B. ClF4+, and SF4 C. ClF4+ and PO43D. XeF4 and ClF4+ E. XeF4, SF4, and PO43-
44. What are the O–Cl–O bond angles in ClO4-? A. 90 B. 109.5 C. 120 D. 90 and 120 E. 180
45. What are the approximate O–N–O bond angles in a nitrate ion? A. 90 B. 109.5 C. 120 D. 90 and 120 E. 180
46. What are the F–Xe–F bond angles in XeF4? A. 90 and 180 B. 109.5 C. 120 D. 60 and 120 E. 180
47. What are the approximate bond angles in SF4? A. 90 B. 109.5 C. 120 D. 90 and 120 E. 180
48. Which of the following molecules has the smallest bond angle between any two hydrogen atoms? A. CH4 B. H2O C. BH3 D. PH3 E. SiH4
49. Which of the following compounds has polar covalent bonds: CCl4, Cl2, HCl, and KCl? A. CCl4 only B. Cl2 only C. HCl and KCl D. Cl2 and KCl E. CCl4 and HCl
50. Which of the following molecules are polar: H2S, CO2, NH3, BH3, and CCl4? A. BH3 B. H2S and NH3 C. H2S, CO2, and CCl4 D. CO2, NH3, and CCl4 E. NH3, BH3 and CCl4
51. What is the hybridization of the carbon atoms in benzene, C6H6? A. sp B. sp2 C. sp3 D. sp3d E. sp3d2
52. What is the hybridization of the phosphorus atom in PCl3? A. sp B. sp2 C. sp3 D. sp3d E. sp3d2
53. What is the hybridization of the central atom in I3-? A. sp B. sp2 C. sp3 D. sp3d E. sp3d2
54. What is the hybridization of the central atom in SO2? A. sp B. sp2 C. sp3 D. sp3d E. sp3d2
55. What hybridization change does the carbon atom undergo in the combustion of methane? CH4(g) + 2O2(g) ® CO2(g) + 2H2O(g) A. sp ® sp2 B. sp2 ® sp3 C. sp3 ® sp D. sp2 ® sp E. none
56. How many sigma and pi bonds are present in the following molecule?
A. 8 sigma bonds and 1 pi bond B. 8 sigma bonds and 2 pi bonds C. 10 sigma bonds and 2 pi bonds D. 11 sigma bonds and 2 pi bonds E. 11 sigma bonds and 1 pi bond
57. How many sigma and pi bonds are present in H2CO? A. 1 sigma bond and 3 pi bonds B. 2 sigma bonds and 2 pi bonds C. 2 sigma bonds and 1 pi bond D. 3 sigma bonds and 1 pi bond E. none of the above
58. G. N. Lewis is well known for multiple contributions to the study of chemistry. Which of these below is attributed to Lewis? A. covalent bonding B. ionic bonding C. VSEPR D. hybridization E. paramagnetism
59. This gas changed the perception of the reactivity of this group. A. oxygen B. xenon C. helium D. hydrogen E. chlorine
60. Write the singly bonded Lewis dot structure for BF3. Which of the following statements best describes this structure? A. It obeys the octet rule on all atoms. B. It has less than an octet on at least one atom. C. It has a lone pair of electrons on the boron atom. D. It has less than an octet of electrons on all atoms. E. It exceeds the octet rule.
61. Write the singly bonded Lewis dot structure for SF6. Which of the following statements best describes this structure? A. It obeys the octet rule on all atoms. B. It has less than an octet on at least one atom. C. It has a lone pair of electrons on the sulfur atom. D. It has less than an octet of electrons on all atoms. E. It exceeds the octet rule.
Chapter 8--Thermochemistry 1. Which of the following is/are state properties? 1. 2. 3.
volume enthalpy heat flow
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 2. Which of the following statements is/are CORRECT? 1. 2. 3.
A system is defined as an object or collection of objects being studied. Surroundings are defined as the entire universe, excluding the system. In an endothermic reaction, heat is transferred from the system to the surroundings.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 3. Which of the following processes is/are endothermic? 1. 2. 3.
the combustion of hydrogen the condensation of water the evaporation of isopropyl alcohol
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3
4. All of the following statements are true EXCEPT A. the value of q is negative in an endothermic process. B. heat flows from the system into the surroundings in an exothermic process. C. the value of q is positive when heat flows into a system from the surroundings. D. enthalpy is a state property. E. in an endothermic process heat flows from the surroundings into the system.
5. How many joules are equivalent to 1.50 ´ 103 calories? A. 1.50 ´ 106 J B. 3.59 ´ 102 J C. 6.28 ´ 103 J D. 3.59 ´ 105 J E. 6.28 ´ 106 J
6. Specific heat (c) is defined as A. the quantity of heat needed to change the temperature of 1.00 g of a substance by 1C. B. the quantity of heat needed to change the temperature of 1.00 g of a substance by 4.184C. C. the capacity of a substance to gain or lose 1.00 J of energy in the form of heat. D. the temperature change undergone when 1.00 g of a substance absorbs 4.184 J. E. the maximum amount of energy in the form of heat that 1.00 g of a substance may absorb without decomposing.
7. If the same amount of energy in the form of heat is added to 5.00 g samples of each of the metals below, which metal will undergo the largest temperature change?
Metal Ag Al Cu Fe Hg
A. Ag B. Al C. Cu D. Fe E. Hg
Specific Heat (J/gC) 0.235 0.897 0.385 0.449 0.140
8. Aluminum has a specific heat of 0.902 J/gC. How many joules of heat are required to change the temperature of 8.50 grams of aluminum from 25.0C to 93.4C? A. 7.67 J B. 71.7 J C. 192 J D. 483 J E. 524 J
9. If 50.0 g of benzene, C6H6, at 25.0C absorbs 2.71 kJ of energy in the form of heat, what is the final temperature of the benzene? The specific heat of benzene is 1.72 J/gC. A. 25.0C B. 31.5C C. 56.5C D. 32.3C E. 57.3C
10. Copper has a specific heat of 0.382 J/gC. The temperature of an unknown mass of copper increases by 4.50C when it absorbs 3.97 J of heat. What is the mass of the copper? A. 2.31 g B. 4.90 g C. 6.82 g D. 8.85 g E. 46.8 g
11. The heat capacity of 5.00 grams of iron is 2.23 J/C. How much heat is required to change the temperature of the iron by 50.0C? A. 0.223 J B. 22.3 J C. 112 J D. 558 J E. 612 J
12. If 495 J is required to change the temperature of 12.7 g of sodium chloride from 75.0C to 135C, what is the specific heat of sodium chloride? A. 0.866 J/gC B. 2.60 J/gC C. 0.650 J/gC D. 1.15 J/gC E. 2.83 ´ 105 J/gC
13. Water has a specific heat of 4.18 J/gC. If 35.0 g of water at 98.8C loses 4.94 kJ of heat, what is the final temperature of the water? A. 32.0C B. 46.2C C. 47.2C D. 57.2C E. 65.0C
14. When 66.0 g of an unknown metal at 28.5C is placed in 83.0 g H2O at 78.5C, the water temperature decreases to 75.9C. What is the specific heat capacity of the metal? The specific heat capacity of water is 4.184 J/gC. A. 0.055 J/gC B. 0.29 J/gC C. 0.69 J/gC D. 0.18 J/gC E. 2.6 J/gC
15. If 35.0 g H2O at 22.7C is combined with 65.0 g H2O at 87.5C, what is the final temperature of the mixture? The specific heat capacity of water is 4.184 J/gC. A. 25.1C B. 45.4C C. 50.8C D. 64.8C E. 48.9C
16. A coffee-cup calorimeter contains 50.0 g of water at 60.51C. A 12.4 g piece of graphite at 24.21C is placed in the calorimeter. The final temperature of the water and the carbon is 59.02C. Calculate the specific heat of carbon. The specific heat of water is 4.18 J/gC. A. 0.328 J/gC B. 0.639 J/gC C. 0.692 J/gC D. 0.721 J/gC E. 1.39 J/gC
17. A coffee-cup calorimeter contains 10.0 g of water at 59.00C. If 3.00 g gold at 15.20C is placed in the calorimeter, what is the final temperature of the water in the calorimeter? The specific heat of water is 4.18 J/gC; the specific heat of gold is 0.128 J/gC. A. 55.37C B. 58.60C C. 59.40C D. 60.80C E. 64.19C
18. When 10.0 g KOH is dissolved in 100.0 g of water in a coffee-cup calorimeter, the temperature rises from 25.18 C to 47.53 C. What is the enthalpy change per gram of KOH dissolved in the water? Assume that the solution has a specific heat capacity of 4.18 J/gK. A. –116 J/g B. –934 J/g C. –1.03 ´ 103 J/g D. –2.19 ´ 103 J/g E. –1.03 ´ 104 J/g
19. When 1.75 g of CaCl2 dissolves in 125 g of water in a coffee-cup calorimeter, the temperature increases by 2.44C. What is the heat change per mole of CaCl2 dissolved in water? Assume that all the heat is absorbed by the water (c = 4.18 J/gC). A. -11.3 kJ B. -728 J C. -1.13 kJ D. -80.9 kJ E. -91.9 kJ
20. Commercial cold packs consist of solid ammonium nitrate and water. NH4NO3 absorbs 25.69 kJ of heat per mole dissolved in water. In a coffee-cup calorimeter, 5.60 g NH4NO3 is dissolved in 100.0 g of water at 22.0C. What is the final temperature of the solution? Assume that the solution has a specific heat capacity of 4.18 J/gC. A. 0.0C B. 17.9C C. 11.6C D. –54.8C E. 26.1C
21. A chemical reaction in a bomb calorimeter evolves 3.15 kJ of heat. If the temperature of the calorimeter is raised from 19.19C to 22.03C, what is the heat capacity of the calorimeter? A. 1.11 kJ/C B. 2.01 kJ/C C. 8.95 kJ/C D. 14.5 kJ/C E. 69.4 kJ/C
22. The combustion of benzoic acid, C7H6O2, can be used to determine the heat capacity of a bomb calorimeter. The heat evolved per mole of benzoic acid combusted is 3.09 ´ 103 kJ. If the combustion of 1.000 g of benzoic acid increases the temperature of a calorimeter by 7.53C. What is the heat capacity of the calorimeter? A. 1.36 kJ/C B. 3.36 kJ/C C. 4.10 kJ/C D. 51.3 kJ/C E. 191 kJ/C
23. Acetylene, C2H2, is a gas used in welding. The molar enthalpy of combustion for acetylene is –2599 kJ. A mass of 0.338 g C2H2(g) is combusted in a bomb calorimeter. If the heat capacity of the calorimeter is 5.54 kJ/C, what is the temperature increase of the bomb calorimeter? A. 1.59C B. 6.09C C. 7.01C D. 12.3C E. 18.0C
24. Isooctane is a primary component of gasoline and gives gasoline its octane rating. Burning 1.00 mL of isooctane (d = 0.688g/mL) releases 33.0 kJ of heat. When 10.0 mL of isooctane is burned in a bomb calorimeter, the temperature in the bomb increases from 23.2C to 66.5C. What is the heat capacity of the bomb calorimeter? A. 11.2 kJ/C B. -3.68 kJ/C C. -7.62 kJ/C D. 3.68 kJ/C E. 7.62 kJ/C
25. Which of the following statements is/are CORRECT? 1. 2. 3.
At constant pressure the heat flow for a reaction equals the change in enthalpy. DH for a reaction is equal in magnitude but opposite in sign to DH for the reverse reaction. Enthalpy is a state function.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 26. Iron oxide reacts with aluminum in an exothermic reaction. Fe2O3(s) + 2Al(s) ® 2Fe(s) + Al2O3(s) The reaction of 5.00 g Fe2O3 with excess Al evolves 26.6 kJ of energy in the form of heat. Calculate the enthalpy change per mole of Fe2O3 reacted. A. –5.32 kJ B. –1.33 ´ 102 kJ C. –2.12 ´ 104 kJ D. –2.12 ´ 102 kJ E. –8.50 ´ 102 kJ
27. Methane, CH4, reacts with oxygen to produce carbon dioxide, water, and heat. CH4(g) + 2O2(g) ® CO2(g) + 2H2O(l)
DH = -890.3 kJ
What is the value of DH if 5.00 g of CH4 is combusted?
A. -157 kJ B. -277 kJ C. -445 kJ D. -714 kJ E. -1.43 ´ 104 kJ 28. If 25.0 g H2O at 85.0C is mixed in a coffee-cup calorimeter with 15.0 g H2O at 20.0C, what is the final temperature of the mixture? The specific heat of water is 4.18 J/gC. A. 45.6C B. 52.5C C. 60.6C D. 68.4C E. 183C
29. 10.0 g of ice at 0.00C is mixed with 25.0 g of water at 35.00C in a coffee-cup calorimeter. What is the final temperature of the mixture? The specific heat of water is 4.18 J/gC; the heat of fusion of water is 333 J/g. A. 0.00C B. 2.24C C. 5.22C D. 25.0C E. 47.8C
30. 30.0 g H2O at an unknown temperature is mixed with 27.0 g of water at 15.8C in a coffee-cup calorimeter. If the final temperature of the mixture is 29.1C, what is the initial temperature of the water? A. 23.7C B. 31.1C C. 39.7C D. 41.1C E. 46.7C
31. Determine the heat of reaction for the decomposition of one mole of benzene to acetylene, C6H6(l) ® 3C2H2(g) given the following thermochemical equations: 2C6H6(l) + 15O2(g) ® 12CO 2(g) + 6H2O(g) 2C2H2(g) + 5O2(g) ® 4CO2(g) + 2H2O(g)
DH = -6271 kJ DH = -2511 kJ
A. 631 kJ B. 1262 kJ C. 3760 kJ D. 6902 kJ E. 8782 kJ 32. Determine the heat of evaporation of carbon disulfide, CS2(l) ® CS2(g) given the enthalpies of reaction below. C(s) + 2S(s) ® CS2(l) C(s) + 2S(s) ® CS2(g)
A. –206.1 kJ B. –27.3 kJ C. +27.3 kJ D. +206.1 kJ E. +1.31 kJ
DH = +89.4 kJ DH = +116.7 kJ
33. Determine the heat of reaction for the combustion of sulfur dioxide, 2SO2(g) + O2(g) ® 2SO3(g) given the following thermochemical equations: S8(s) + 8O2(g) ® 8SO2(g) S8(s) + 12O2(g) ® 8SO3(g)
DH = -2374.6 kJ DH = -3165.8 kJ
A. -5540.4 kJ B. -1385.1 kJ C. -197.8 kJ D. 251.7 kJ E. 791.2 kJ 34. Determine the heat of formation of sulfuric acid, H2SO4(l), from the following thermochemical equations: S8(s) + 8O2(g) ® 8SO2(g) S8(s) + 12O2(g) ® 8SO3(g) H2O(l) + SO3(g) ® H2SO4(l) 2H2(g) + O2(g) ® 2H2O(l)
DH = -2374.6 kJ DH = -3165.8 kJ DH = -132.4 kJ DH = -571.7 kJ
A. -814.0 kJ B. -1099.9 kJ C. -2164.7 kJ D. -3869.9 kJ E. -6244.5 kJ 35. Determine DH for the following reaction, 2NH3(g) + 5/2O2(g) ® 2NO(g) + 3H2O(g) given the thermochemical equations below. N2(g) + O2(g) ® 2NO(g) N2(g) + 3H2(g) ® 2NH3(g) 2H2(g) + O2(g) ® 2H2O(g)
A. –1178.2 kJ B. –452.8 kJ C. –394.6 kJ D. –211.0 kJ E. +1178.2 kJ
DH = +180.8 kJ DH = –91.8 kJ DH = –483.6 kJ
36. Determine the heat of reaction for the oxidation of iron, 2Fe(s) + O2(g) ® Fe2O3(s) given the following thermochemical equations: 2Fe(s) + 6H2O(l) ® 2Fe(OH)3(s) + 3H2(g) Fe2O3(s) + 3H2O(l) ® 2Fe(OH)3(s) 2H2(g) + O2(g) ® 2H2O(l)
DH = 321.8 kJ DH = 288.6 kJ DH = -571.7 kJ
A. -1681.9 kJ B. -1143.1 kJ C. -824.4 kJ D. 33.2 kJ E. 38.7 kJ 37. Determine the heat of formation of calcium carbonate from the thermochemical equations given below. Ca(OH)2(s) ® CaO(s) + H2O(l) Ca(OH)2(s) + CO2(g) ® CaCO3(s) + H2O(l) C(s) + O2(g) ® CO2(g) 2Ca(s) + O2(g) ® 2CaO(s)
DH = 65.2 kJ DH = -113.2 kJ DH = -393.5 kJ DH = -1270.2 kJ
A. -178.4 kJ B. -493.2 kJ C. -828.7 kJ D. -980.6 kJ E. -1207.0 kJ 38. All of the following statements are true EXCEPT A. Hess' law states that DH for an overall reaction is the sum of the DH values for the individual equations. B. the molar enthalpy of formation of a compound is equal to the enthalpy change when one mole of the compound is formed from elements. C. a reaction with a negative enthalpy is exothermic. D. the enthalpy of formation of an element in its most stable state is equal to zero. E. the sum of the enthalpies of formation of the products in a chemical reaction is defined as the enthalpy of reaction.
39. Determine the heat of reaction for the combustion of ammonia, 4NH3(g) + 7O2(g) ® 4NO2(g) + 6H2O(l) using molar enthalpies of formation.
molecule NH3(g) NO2(g) H2O(l)
DfH (kJ/mol) –45.9 +33.1 –285.8
A. +30.24 kJ B. –206.9 kJ C. –298.6 kJ D. –1398.8 kJ E. –1663.6 kJ 40. Using molar enthalpies of formation, determine the heat of reaction for the combustion of 1.0000 mole of methanol. DHf (kJ/mol)
2CH3OH(l) -238.7
+
3O2(g)
®
2CO2(g) -393.5
+
4H2O(l) -285.8
A. -2407.6 kJ B. -1452.8 kJ C. -918.0 kJ D. -726.4 kJ E. -148.9 kJ 41. The molar enthalpies of formation for H2O(l) and H2O(g) are -285.8 kJ and -241.8 kJ, respectively. How much heat is released when 25 g of water condenses from the gas to the liquid phase? A. -2.4 kJ B. -32 kJ C. -61 kJ D. -88 kJ E. -1100 kJ
42. If not handled carefully, ammonium perchlorate can decompose violently according to the thermochemical equation below. 2NH4ClO4(s) ® N2(g) + Cl2(g) + 2O2(g) + 4H2O(g)
DH = -375.6 kJ
The enthalpy of formation of H2O(g) is -241.8 kJ. Calculate the enthalpy of formation of ammonium perchlorate.
A. -295.8 kJ B. -156.4 kJ C. 66.9 kJ D. 133.8 kJ E. 561.1 kJ 43. The thermite reaction is an exothermic process that yields iron metal as a product. DHf (kJ/mol)
Fe2O3(s) -824.2
+
2Al(s)
®
2Fe(s)
+
Al2O3(s) -1675.7
Using enthalpies of formation, determine the heat released when 5.00 g of Fe2O3(s) reacts with excess Al.
A. -26.7 kJ B. -171 kJ C. -243 kJ D. -850.8 kJ E. -2720 kJ 44. The standard molar enthalpy of formation of NH3(g) is –45.9 kJ/mol. What is the enthalpy change if 9.51 g N2(g) and 1.96 g H2(g) react to produce NH3(g)? A. –10.3 kJ B. –20.7 kJ C. –29.8 kJ D. –43.7 kJ E. –65.6 kJ
45. The enthalpy of formation of H2O(l) is -285.8 kJ/mol. What is the enthalpy change if 1.62 g H2(g) reacts with 9.21 g O2(g) to form H2O(l)? A. -229 kJ B. -165 kJ C. -107 kJ D. -45.5 kJ E. -31.0 kJ
46. Carbon and oxygen react to give carbon dioxide. The reaction of 4.49 g C(s) with 9.21 g O2(g) releases -113.2 kJ of heat. What is the enthalpy of formation of CO2(g)? A. -393 kJ B. -303 kJ C. -285.8 kJ D. -171 kJ E. -113.2 kJ
47. Which of the following chemical equations does not correspond to a standard molar enthalpy of formation? A. Mg(s) + C(s) + 3/2O2(g) ® MgCO3(s) B. C(s) + 1/2O2(g) ® CO(g) C. N2(g) + O2(g) ® 2NO(g) D. N2(g) + 2O2(g) ® N2O4(g) E. H2(g) + 1/2O2(g) ® H2O(l)
48. In which of the reactions below is DH an enthalpy of formation? A. S8(s) + 8O2(g) ® 8SO2(g) B. CaO(s) + H2O(l) ® Ca(OH)2(s) C. H2(g) + Cl2(g) + 4O2(g) ® 2HClO4(l) D. Cu(s) + 1/2I2(s) ® CuI(s) E. CH4(g) + 2O2(g) ® CO2(g) + 2H2O(g)
49. Calculate the enthalpy change for the following reaction, Br2(g) + 3F2(g) ® 2BrF3(g) given the bond enthalpies of the reactants and products.
Bond Br-Br F-F Br-F
Bond Enthalpy (kJ) 193 155 249
A. –836 kJ B. –89 kJ C. +89 kJ D. +99 kJ E. +836 kJ 50. Using the bond enthalpies tabulated below, calculate the enthalpy of reaction for the combustion of hydrogen. 2H2(g) + O2(g) ® 2H2O(g)
bond H-H O-H O-O O=O
A. -846 kJ B. -486 kJ C. -348 kJ D. -243 kJ E. 888 kJ
DH (kJ) 436 464 138 498
51. Using the bond enthalpies tabulated below, calculate the enthalpy of reaction for the combustion of ethane. 2C2H6(g) + 7O2(g) ® 4CO2(g) + 6H2O(g)
bond H-H C-H C-C C-O
DH (kJ) 436 414 347 351
bond C=O O-O O=O O-H
DH (kJ) 715 138 498 464
A. -2834 kJ B. -2140 kJ C. -1174 kJ D. -644 kJ E. -322 kJ 52. Which of the following statements is/are CORRECT? 1. 2. 3.
The formation of a bond is an endothermic process. A reaction is expected to be exothermic if the bonds of the products are stronger than those of the reactants. The bond energy of a double bond is exactly twice that of a single bond..
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 53. The first law of thermodynamics states that A. the amount of work done on a system is independent of the pathway. B. the heat flow in or out of a system is independent of the pathway. C. the total energy change of a system is equal to the sum of the heat absorbed and the work done on the system. D. both answers a and c E. both answers b and c
54. Calculate DE of a gas for a process in which the gas absorbs 25 J of heat and does 7 J of work by expanding. A. -32 J B. -18 J C. +18 J D. +32 J E. +180 J
55. Calculate DE of a gas for a process in which the gas evolves 8 J of heat and does 16 J of work by expanding. A. +24 J B. -2 J C. +8 J D. -24 J E. -8 J
56. Calculate DE of a gas for a process in which the gas evolves 24 J of heat and has 9 J of work done on it. A. -33 J B. -15 J C. +15 J D. +33 J E. -220 J
57. All of the following statements are correct EXCEPT A. in any chemical reaction the change in energies of the system and the surroundings are equal in magnitude but opposite in sign. B. an expanding gas does work on the surroundings, thus w has a negative value. C. DE = q + w (where q equals heat and w equals work). D. if a reaction occurs at constant volume, q = DE. E. if a reaction occurs at constant pressure, w = DE.
58. You decide to go on a diet to lose 5 lbs (1 lb = 454 g). Fatty tissue consists of about 85% fat and 15% water. In order to lose that much weight how many Calories (kcal) must be eliminated from your normal intake? Fat contains 9.0 kcal/g. A. 17400 kcal B. 20600 kcal C. 3100 kcal D. 3500 kcal E. 400 kcal
59. How many kJ are equal to 3.27 Latm of work? A. 0.331 kJ B. 3.23 kJ C. 331 kJ D. 0.0323 kJ E. 33.0 kJ
Chapter 8--Thermochemistry Key
1. Which of the following is/are state properties? 1. 2. 3.
volume enthalpy heat flow
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 2. Which of the following statements is/are CORRECT? 1. 2. 3.
A system is defined as an object or collection of objects being studied. Surroundings are defined as the entire universe, excluding the system. In an endothermic reaction, heat is transferred from the system to the surroundings.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 3. Which of the following processes is/are endothermic? 1. 2. 3.
the combustion of hydrogen the condensation of water the evaporation of isopropyl alcohol
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3
4. All of the following statements are true EXCEPT A. the value of q is negative in an endothermic process. B. heat flows from the system into the surroundings in an exothermic process. C. the value of q is positive when heat flows into a system from the surroundings. D. enthalpy is a state property. E. in an endothermic process heat flows from the surroundings into the system.
5. How many joules are equivalent to 1.50 ´ 103 calories? A. 1.50 ´ 106 J B. 3.59 ´ 102 J C. 6.28 ´ 103 J D. 3.59 ´ 105 J E. 6.28 ´ 106 J
6. Specific heat (c) is defined as A. the quantity of heat needed to change the temperature of 1.00 g of a substance by 1C. B. the quantity of heat needed to change the temperature of 1.00 g of a substance by 4.184C. C. the capacity of a substance to gain or lose 1.00 J of energy in the form of heat. D. the temperature change undergone when 1.00 g of a substance absorbs 4.184 J. E. the maximum amount of energy in the form of heat that 1.00 g of a substance may absorb without decomposing.
7. If the same amount of energy in the form of heat is added to 5.00 g samples of each of the metals below, which metal will undergo the largest temperature change?
Metal Ag Al Cu Fe Hg
A. Ag B. Al C. Cu D. Fe E. Hg
Specific Heat (J/gC) 0.235 0.897 0.385 0.449 0.140
8. Aluminum has a specific heat of 0.902 J/gC. How many joules of heat are required to change the temperature of 8.50 grams of aluminum from 25.0C to 93.4C? A. 7.67 J B. 71.7 J C. 192 J D. 483 J E. 524 J
9. If 50.0 g of benzene, C6H6, at 25.0C absorbs 2.71 kJ of energy in the form of heat, what is the final temperature of the benzene? The specific heat of benzene is 1.72 J/gC. A. 25.0C B. 31.5C C. 56.5C D. 32.3C E. 57.3C
10. Copper has a specific heat of 0.382 J/gC. The temperature of an unknown mass of copper increases by 4.50C when it absorbs 3.97 J of heat. What is the mass of the copper? A. 2.31 g B. 4.90 g C. 6.82 g D. 8.85 g E. 46.8 g
11. The heat capacity of 5.00 grams of iron is 2.23 J/C. How much heat is required to change the temperature of the iron by 50.0C? A. 0.223 J B. 22.3 J C. 112 J D. 558 J E. 612 J
12. If 495 J is required to change the temperature of 12.7 g of sodium chloride from 75.0C to 135C, what is the specific heat of sodium chloride? A. 0.866 J/gC B. 2.60 J/gC C. 0.650 J/gC D. 1.15 J/gC E. 2.83 ´ 105 J/gC
13. Water has a specific heat of 4.18 J/gC. If 35.0 g of water at 98.8C loses 4.94 kJ of heat, what is the final temperature of the water? A. 32.0C B. 46.2C C. 47.2C D. 57.2C E. 65.0C
14. When 66.0 g of an unknown metal at 28.5C is placed in 83.0 g H2O at 78.5C, the water temperature decreases to 75.9C. What is the specific heat capacity of the metal? The specific heat capacity of water is 4.184 J/gC. A. 0.055 J/gC B. 0.29 J/gC C. 0.69 J/gC D. 0.18 J/gC E. 2.6 J/gC
15. If 35.0 g H2O at 22.7C is combined with 65.0 g H2O at 87.5C, what is the final temperature of the mixture? The specific heat capacity of water is 4.184 J/gC. A. 25.1C B. 45.4C C. 50.8C D. 64.8C E. 48.9C
16. A coffee-cup calorimeter contains 50.0 g of water at 60.51C. A 12.4 g piece of graphite at 24.21C is placed in the calorimeter. The final temperature of the water and the carbon is 59.02C. Calculate the specific heat of carbon. The specific heat of water is 4.18 J/gC. A. 0.328 J/gC B. 0.639 J/gC C. 0.692 J/gC D. 0.721 J/gC E. 1.39 J/gC
17. A coffee-cup calorimeter contains 10.0 g of water at 59.00C. If 3.00 g gold at 15.20C is placed in the calorimeter, what is the final temperature of the water in the calorimeter? The specific heat of water is 4.18 J/gC; the specific heat of gold is 0.128 J/gC. A. 55.37C B. 58.60C C. 59.40C D. 60.80C E. 64.19C
18. When 10.0 g KOH is dissolved in 100.0 g of water in a coffee-cup calorimeter, the temperature rises from 25.18 C to 47.53 C. What is the enthalpy change per gram of KOH dissolved in the water? Assume that the solution has a specific heat capacity of 4.18 J/gK. A. –116 J/g B. –934 J/g C. –1.03 ´ 103 J/g D. –2.19 ´ 103 J/g E. –1.03 ´ 104 J/g
19. When 1.75 g of CaCl2 dissolves in 125 g of water in a coffee-cup calorimeter, the temperature increases by 2.44C. What is the heat change per mole of CaCl2 dissolved in water? Assume that all the heat is absorbed by the water (c = 4.18 J/gC). A. -11.3 kJ B. -728 J C. -1.13 kJ D. -80.9 kJ E. -91.9 kJ
20. Commercial cold packs consist of solid ammonium nitrate and water. NH4NO3 absorbs 25.69 kJ of heat per mole dissolved in water. In a coffee-cup calorimeter, 5.60 g NH4NO3 is dissolved in 100.0 g of water at 22.0C. What is the final temperature of the solution? Assume that the solution has a specific heat capacity of 4.18 J/gC. A. 0.0C B. 17.9C C. 11.6C D. –54.8C E. 26.1C
21. A chemical reaction in a bomb calorimeter evolves 3.15 kJ of heat. If the temperature of the calorimeter is raised from 19.19C to 22.03C, what is the heat capacity of the calorimeter? A. 1.11 kJ/C B. 2.01 kJ/C C. 8.95 kJ/C D. 14.5 kJ/C E. 69.4 kJ/C
22. The combustion of benzoic acid, C7H6O2, can be used to determine the heat capacity of a bomb calorimeter. The heat evolved per mole of benzoic acid combusted is 3.09 ´ 103 kJ. If the combustion of 1.000 g of benzoic acid increases the temperature of a calorimeter by 7.53C. What is the heat capacity of the calorimeter? A. 1.36 kJ/C B. 3.36 kJ/C C. 4.10 kJ/C D. 51.3 kJ/C E. 191 kJ/C
23. Acetylene, C2H2, is a gas used in welding. The molar enthalpy of combustion for acetylene is –2599 kJ. A mass of 0.338 g C2H2(g) is combusted in a bomb calorimeter. If the heat capacity of the calorimeter is 5.54 kJ/C, what is the temperature increase of the bomb calorimeter? A. 1.59C B. 6.09C C. 7.01C D. 12.3C E. 18.0C
24. Isooctane is a primary component of gasoline and gives gasoline its octane rating. Burning 1.00 mL of isooctane (d = 0.688g/mL) releases 33.0 kJ of heat. When 10.0 mL of isooctane is burned in a bomb calorimeter, the temperature in the bomb increases from 23.2C to 66.5C. What is the heat capacity of the bomb calorimeter? A. 11.2 kJ/C B. -3.68 kJ/C C. -7.62 kJ/C D. 3.68 kJ/C E. 7.62 kJ/C
25. Which of the following statements is/are CORRECT? 1. 2. 3.
At constant pressure the heat flow for a reaction equals the change in enthalpy. DH for a reaction is equal in magnitude but opposite in sign to DH for the reverse reaction. Enthalpy is a state function.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 26. Iron oxide reacts with aluminum in an exothermic reaction. Fe2O3(s) + 2Al(s) ® 2Fe(s) + Al2O3(s) The reaction of 5.00 g Fe2O3 with excess Al evolves 26.6 kJ of energy in the form of heat. Calculate the enthalpy change per mole of Fe2O3 reacted. A. –5.32 kJ B. –1.33 ´ 102 kJ C. –2.12 ´ 104 kJ D. –2.12 ´ 102 kJ E. –8.50 ´ 102 kJ
27. Methane, CH4, reacts with oxygen to produce carbon dioxide, water, and heat. CH4(g) + 2O2(g) ® CO2(g) + 2H2O(l)
DH = -890.3 kJ
What is the value of DH if 5.00 g of CH4 is combusted?
A. -157 kJ B. -277 kJ C. -445 kJ D. -714 kJ E. -1.43 ´ 104 kJ 28. If 25.0 g H2O at 85.0C is mixed in a coffee-cup calorimeter with 15.0 g H2O at 20.0C, what is the final temperature of the mixture? The specific heat of water is 4.18 J/gC. A. 45.6C B. 52.5C C. 60.6C D. 68.4C E. 183C
29. 10.0 g of ice at 0.00C is mixed with 25.0 g of water at 35.00C in a coffee-cup calorimeter. What is the final temperature of the mixture? The specific heat of water is 4.18 J/gC; the heat of fusion of water is 333 J/g. A. 0.00C B. 2.24C C. 5.22C D. 25.0C E. 47.8C
30. 30.0 g H2O at an unknown temperature is mixed with 27.0 g of water at 15.8C in a coffee-cup calorimeter. If the final temperature of the mixture is 29.1C, what is the initial temperature of the water? A. 23.7C B. 31.1C C. 39.7C D. 41.1C E. 46.7C
31. Determine the heat of reaction for the decomposition of one mole of benzene to acetylene, C6H6(l) ® 3C2H2(g) given the following thermochemical equations: 2C6H6(l) + 15O2(g) ® 12CO 2(g) + 6H2O(g) 2C2H2(g) + 5O2(g) ® 4CO2(g) + 2H2O(g)
DH = -6271 kJ DH = -2511 kJ
A. 631 kJ B. 1262 kJ C. 3760 kJ D. 6902 kJ E. 8782 kJ 32. Determine the heat of evaporation of carbon disulfide, CS2(l) ® CS2(g) given the enthalpies of reaction below. C(s) + 2S(s) ® CS2(l) C(s) + 2S(s) ® CS2(g)
A. –206.1 kJ B. –27.3 kJ C. +27.3 kJ D. +206.1 kJ E. +1.31 kJ
DH = +89.4 kJ DH = +116.7 kJ
33. Determine the heat of reaction for the combustion of sulfur dioxide, 2SO2(g) + O2(g) ® 2SO3(g) given the following thermochemical equations: S8(s) + 8O2(g) ® 8SO2(g) S8(s) + 12O2(g) ® 8SO3(g)
DH = -2374.6 kJ DH = -3165.8 kJ
A. -5540.4 kJ B. -1385.1 kJ C. -197.8 kJ D. 251.7 kJ E. 791.2 kJ 34. Determine the heat of formation of sulfuric acid, H2SO4(l), from the following thermochemical equations: S8(s) + 8O2(g) ® 8SO2(g) S8(s) + 12O2(g) ® 8SO3(g) H2O(l) + SO3(g) ® H2SO4(l) 2H2(g) + O2(g) ® 2H2O(l)
DH = -2374.6 kJ DH = -3165.8 kJ DH = -132.4 kJ DH = -571.7 kJ
A. -814.0 kJ B. -1099.9 kJ C. -2164.7 kJ D. -3869.9 kJ E. -6244.5 kJ 35. Determine DH for the following reaction, 2NH3(g) + 5/2O2(g) ® 2NO(g) + 3H2O(g) given the thermochemical equations below. N2(g) + O2(g) ® 2NO(g) N2(g) + 3H2(g) ® 2NH3(g) 2H2(g) + O2(g) ® 2H2O(g)
A. –1178.2 kJ B. –452.8 kJ C. –394.6 kJ D. –211.0 kJ E. +1178.2 kJ
DH = +180.8 kJ DH = –91.8 kJ DH = –483.6 kJ
36. Determine the heat of reaction for the oxidation of iron, 2Fe(s) + O2(g) ® Fe2O3(s) given the following thermochemical equations: 2Fe(s) + 6H2O(l) ® 2Fe(OH)3(s) + 3H2(g) Fe2O3(s) + 3H2O(l) ® 2Fe(OH)3(s) 2H2(g) + O2(g) ® 2H2O(l)
DH = 321.8 kJ DH = 288.6 kJ DH = -571.7 kJ
A. -1681.9 kJ B. -1143.1 kJ C. -824.4 kJ D. 33.2 kJ E. 38.7 kJ 37. Determine the heat of formation of calcium carbonate from the thermochemical equations given below. Ca(OH)2(s) ® CaO(s) + H2O(l) Ca(OH)2(s) + CO2(g) ® CaCO3(s) + H2O(l) C(s) + O2(g) ® CO2(g) 2Ca(s) + O2(g) ® 2CaO(s)
DH = 65.2 kJ DH = -113.2 kJ DH = -393.5 kJ DH = -1270.2 kJ
A. -178.4 kJ B. -493.2 kJ C. -828.7 kJ D. -980.6 kJ E. -1207.0 kJ 38. All of the following statements are true EXCEPT A. Hess' law states that DH for an overall reaction is the sum of the DH values for the individual equations. B. the molar enthalpy of formation of a compound is equal to the enthalpy change when one mole of the compound is formed from elements. C. a reaction with a negative enthalpy is exothermic. D. the enthalpy of formation of an element in its most stable state is equal to zero. E. the sum of the enthalpies of formation of the products in a chemical reaction is defined as the enthalpy of reaction.
39. Determine the heat of reaction for the combustion of ammonia, 4NH3(g) + 7O2(g) ® 4NO2(g) + 6H2O(l) using molar enthalpies of formation.
molecule NH3(g) NO2(g) H2O(l)
DfH (kJ/mol) –45.9 +33.1 –285.8
A. +30.24 kJ B. –206.9 kJ C. –298.6 kJ D. –1398.8 kJ E. –1663.6 kJ 40. Using molar enthalpies of formation, determine the heat of reaction for the combustion of 1.0000 mole of methanol. DHf (kJ/mol)
2CH3OH(l) -238.7
+
3O2(g)
®
2CO2(g) -393.5
+
4H2O(l) -285.8
A. -2407.6 kJ B. -1452.8 kJ C. -918.0 kJ D. -726.4 kJ E. -148.9 kJ 41. The molar enthalpies of formation for H2O(l) and H2O(g) are -285.8 kJ and -241.8 kJ, respectively. How much heat is released when 25 g of water condenses from the gas to the liquid phase? A. -2.4 kJ B. -32 kJ C. -61 kJ D. -88 kJ E. -1100 kJ
42. If not handled carefully, ammonium perchlorate can decompose violently according to the thermochemical equation below. 2NH4ClO4(s) ® N2(g) + Cl2(g) + 2O2(g) + 4H2O(g)
DH = -375.6 kJ
The enthalpy of formation of H2O(g) is -241.8 kJ. Calculate the enthalpy of formation of ammonium perchlorate.
A. -295.8 kJ B. -156.4 kJ C. 66.9 kJ D. 133.8 kJ E. 561.1 kJ 43. The thermite reaction is an exothermic process that yields iron metal as a product. DHf (kJ/mol)
Fe2O3(s) -824.2
+
2Al(s)
®
2Fe(s)
+
Al2O3(s) -1675.7
Using enthalpies of formation, determine the heat released when 5.00 g of Fe2O3(s) reacts with excess Al.
A. -26.7 kJ B. -171 kJ C. -243 kJ D. -850.8 kJ E. -2720 kJ 44. The standard molar enthalpy of formation of NH3(g) is –45.9 kJ/mol. What is the enthalpy change if 9.51 g N2(g) and 1.96 g H2(g) react to produce NH3(g)? A. –10.3 kJ B. –20.7 kJ C. –29.8 kJ D. –43.7 kJ E. –65.6 kJ
45. The enthalpy of formation of H2O(l) is -285.8 kJ/mol. What is the enthalpy change if 1.62 g H2(g) reacts with 9.21 g O2(g) to form H2O(l)? A. -229 kJ B. -165 kJ C. -107 kJ D. -45.5 kJ E. -31.0 kJ
46. Carbon and oxygen react to give carbon dioxide. The reaction of 4.49 g C(s) with 9.21 g O2(g) releases -113.2 kJ of heat. What is the enthalpy of formation of CO2(g)? A. -393 kJ B. -303 kJ C. -285.8 kJ D. -171 kJ E. -113.2 kJ
47. Which of the following chemical equations does not correspond to a standard molar enthalpy of formation? A. Mg(s) + C(s) + 3/2O2(g) ® MgCO3(s) B. C(s) + 1/2O2(g) ® CO(g) C. N2(g) + O2(g) ® 2NO(g) D. N2(g) + 2O2(g) ® N2O4(g) E. H2(g) + 1/2O2(g) ® H2O(l)
48. In which of the reactions below is DH an enthalpy of formation? A. S8(s) + 8O2(g) ® 8SO2(g) B. CaO(s) + H2O(l) ® Ca(OH)2(s) C. H2(g) + Cl2(g) + 4O2(g) ® 2HClO4(l) D. Cu(s) + 1/2I2(s) ® CuI(s) E. CH4(g) + 2O2(g) ® CO2(g) + 2H2O(g)
49. Calculate the enthalpy change for the following reaction, Br2(g) + 3F2(g) ® 2BrF3(g) given the bond enthalpies of the reactants and products.
Bond Br-Br F-F Br-F
Bond Enthalpy (kJ) 193 155 249
A. –836 kJ B. –89 kJ C. +89 kJ D. +99 kJ E. +836 kJ 50. Using the bond enthalpies tabulated below, calculate the enthalpy of reaction for the combustion of hydrogen. 2H2(g) + O2(g) ® 2H2O(g)
bond H-H O-H O-O O=O
A. -846 kJ B. -486 kJ C. -348 kJ D. -243 kJ E. 888 kJ
DH (kJ) 436 464 138 498
51. Using the bond enthalpies tabulated below, calculate the enthalpy of reaction for the combustion of ethane. 2C2H6(g) + 7O2(g) ® 4CO2(g) + 6H2O(g)
bond H-H C-H C-C C-O
DH (kJ) 436 414 347 351
bond C=O O-O O=O O-H
DH (kJ) 715 138 498 464
A. -2834 kJ B. -2140 kJ C. -1174 kJ D. -644 kJ E. -322 kJ 52. Which of the following statements is/are CORRECT? 1. 2. 3.
The formation of a bond is an endothermic process. A reaction is expected to be exothermic if the bonds of the products are stronger than those of the reactants. The bond energy of a double bond is exactly twice that of a single bond..
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 53. The first law of thermodynamics states that A. the amount of work done on a system is independent of the pathway. B. the heat flow in or out of a system is independent of the pathway. C. the total energy change of a system is equal to the sum of the heat absorbed and the work done on the system. D. both answers a and c E. both answers b and c
54. Calculate DE of a gas for a process in which the gas absorbs 25 J of heat and does 7 J of work by expanding. A. -32 J B. -18 J C. +18 J D. +32 J E. +180 J
55. Calculate DE of a gas for a process in which the gas evolves 8 J of heat and does 16 J of work by expanding. A. +24 J B. -2 J C. +8 J D. -24 J E. -8 J
56. Calculate DE of a gas for a process in which the gas evolves 24 J of heat and has 9 J of work done on it. A. -33 J B. -15 J C. +15 J D. +33 J E. -220 J
57. All of the following statements are correct EXCEPT A. in any chemical reaction the change in energies of the system and the surroundings are equal in magnitude but opposite in sign. B. an expanding gas does work on the surroundings, thus w has a negative value. C. DE = q + w (where q equals heat and w equals work). D. if a reaction occurs at constant volume, q = DE. E. if a reaction occurs at constant pressure, w = DE.
58. You decide to go on a diet to lose 5 lbs (1 lb = 454 g). Fatty tissue consists of about 85% fat and 15% water. In order to lose that much weight how many Calories (kcal) must be eliminated from your normal intake? Fat contains 9.0 kcal/g. A. 17400 kcal B. 20600 kcal C. 3100 kcal D. 3500 kcal E. 400 kcal
59. How many kJ are equal to 3.27 Latm of work? A. 0.331 kJ B. 3.23 kJ C. 331 kJ D. 0.0323 kJ E. 33.0 kJ
Chapter 9--Liquids and Solids 1. Equilibrium has been established between a liquid and its vapor when A. the masses of liquid and vapor are equal. B. all of the liquid has evaporated. C. when the temperatures of the liquid and the vapor are the same. D. evaporation ceases and the concentrations of liquid and vapor remain constant. E. the rate of condensation equals the rate of evaporation.
2. Which of the following statements is/are CORRECT? The vapor pressure of a liquid depends on 1. 2. 3.
the temperature of the liquid. the surface area of the liquid. the volume of the liquid.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 3. Methyl alcohol, CH3OH, has a vapor pressure of 203 mm Hg at 35C. If 5.00 g CH3OH is sealed in a 10.0 L flask, what mass will remain in the liquid phase when equilibrium is established at 35C? Assume any liquid remaining in the flask has a negligible volume. (760 mm Hg = 1 atm, R = 0.0821 Latm/molK) A. 0.00 g B. 1.62 g C. 2.08 g D. 2.27 g E. 4.69 g
4. At 75.0 C, water has an equilibrium vapor pressure of 289.1 mm Hg. If 4.22 g H2O is sealed in an evacuated 5.00 L flask and heated to 75.0 C, what mass of H2O will be found in the gas phase when liquid-vapor equilibrium is established? Assume any liquid remaining in the flask has a negligible volume. (760 mm Hg = 1 atm, R = 0.0821 Latm/molK) A. 0.240 g B. 1.20 g C. 2.64 g D. 3.02 g E. 4.22 g
5. Water has an equilibrium vapor pressure of 23.8 mm Hg at 25C. What mass of water vapor is present, at equilibrium, in a room with dimensions of 7.0 m ´ 8.0 m ´ 2.7 m? (760 mm Hg = 1 atm, R = 0.0821 Latm/molK) A. 0.0050 kg B. 0.029 kg C. 0.28 kg D. 3.5 kg E. 5.0 kg
6. Which of the following statements is true for an ideal gas? A. A plot of ln P vsersus 1/T (in Kelvin) yields a straight line with a slope equal to -DHvap. B. A plot of ln P versus T (in Kelvin) yields a straight line with a slope equal to DHvap. C. A plot of ln P versus 1/T (in Kelvin) yields a straight line with a slope equal to -DHvap/R. D. A plot of P versus 1/T (in Kelvin) yields a straight line with a slope equal to 1/-DHvap. E. A plot of P versus T (in Kelvin) yields a straight line with a slope equal to -DHvap/RT.
7. Which of the following equations is a correct form of the Clausius-Clapeyron equation?
A.
B.
C.
D.
E.
8. Sulfur dioxide has a vapor pressure of 462.7 mm Hg at –21.0C and a vapor pressure of 140.5 mm Hg at –44.0C. What is the molar heat of vaporization of sulfur dioxide? (R = 8.31 J/Kmol) A. 0.398 kJ/mol B. 6.33 kJ/mol C. 14.0 kJ/mol D. 24.9 kJ/mol E. 39.8 kJ/mol
9. Ethanol has a molar heat of vaporization of 42.3 kJ/mol. The compound has a vapor pressure of 1.00 atm at 78.3C. At what temperature is the vapor pressure equal to 0.800 atm? (R = 8.31 J/Kmol) A. –83.8C B. –24.4C C. 62.6C D. 73.0C E. 78.0C
10. Carbon tetrachloride, an organic solvent, has a molar heat of vaporization of 29.82 kJ/mol. If CCl4 has a normal boiling point of 3.50 ´ 102 K, what is its vapor pressure at 273 K? (R = 8.31 J/molK) A. 19.3 mm Hg B. 42.2 mm Hg C. 59.3 mm Hg D. 108 mm Hg E. 359 mm Hg
11. Mount Everest rises to a height of 29,035 ft (8.850 ´ 103 m) above sea level. At this height, the atmospheric pressure is 230 mm Hg. At what temperature does water boil at the summit of Mount Everest? (DHvap for H2O = 40.7 kJ/mole, R = 8.31 J/Kmol) A. -17C B. 31C C. 48C D. 57C E. 69C
12. Sulfur dioxide has an enthalpy of vaporization of 24.9 kJ/mol. At 205 K, SO2 has a vapor pressure of 30.3 mm Hg. What is the normal boiling point temperature of SO2? (R = 8.31 J/Kmol) A. 263 K B. 278 K C. 291 K D. 308 K E. 345 K
13. Freon-113, C2Cl3F3, has an enthalpy of vaporization of 27.04 kJ/mol and a normal boiling point of 48C. At what temperature is the vapor pressure of Freon-113 equal to 76.0 mm Hg? (R = 8.31 J/Kmol) A. -53.1C B. -23.8C C. -11.4C D. 1.44C E. 7.12C
14. Xenon has a molar heat of vaporization of 12.6 kJ/mol and a vapor pressure of 1.00 atm at –108.0C. What is the vapor pressure of xenon at -148.0C? (R = 8.31 J/Kmol) A. 0.053 atm B. 0.73 atm C. 0.93 atm D. 0.99 atm E. 19 atm
15. Naphthalene, a substance present in some mothballs, has a molar heat of vaporization of 49.4 kJ/mol. If the vapor pressure of naphthalene is 0.300 mm Hg at 298 K, what is the temperature at which the vapor pressure is 7.60 ´ 102 mm Hg? (R = 8.31 J/Kmol) A. 424 K B. 491 K C. 501 K D. 521 K E. 611 K
16. The normal boiling point of a liquid is A. 373 K. B. the temperature at which a liquid's vapor pressure equals 1 atm. C. the pressure at which the liquid boils at 373 K. D. dependent upon the volume of the liquid. E. dependent upon the surface area of the liquid..
17. All of the following statements are incorrect EXCEPT A. gas, liquid and solid phases of a substance exist simultaneously at equilibrium at the critical pressure and temperature. B. above the critical pressure, only the solid phase of a pure substance can exist. C. the liquid phase of a pure substance cannot exist above its critical temperature. D. the critical temperature and pressure cannot be produced simultaneously. E. above the critical temperature and pressure, only the liquid phase of a substance exists.
18. Which of the following statements are correct? 1. 2. 3.
All three phases (gas, liquid, and solid) exist at equilibrium at the triple point. Above its critical temperature and pressure, a substance becomes a supercritical fluid. A liquid boils when its vapor pressure equals the pressure above its surface.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 19. The phase diagram for CO2 has a triple point at –56.6 C and 5.19 atm, and a critical point at 31.0 C and 73 atm. The solid and gas phases are in equilibrium at –78.7 C and 1.00 atm. Which of the following statements regarding CO2 is/are CORRECT? 1. 2. 3.
Sublimation occurs if the temperature of the solid phase is increased from –79.0 C to 0.0 C at a constant pressure of 2.5 atm. CO2 is a supercritical fluid at 55 C and 75 atm. At pressures greater than its critical pressure (73 atm), CO2 will not exist as a solid at any temperature.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 20. On the phase diagram below, which point corresponds to conditions where only the solid phase exists?
A. A B. B C. C D. D E. G
21. On the phase diagram below, which point corresponds to conditions where both liquid and gas phases exist?
A. A B. B C. C D. E E. F
22. On the phase diagram below, which point corresponds to conditions where solid, liquid, and gas phases all exist?
A. B B. C C. D D. E E. G
23. If a pure substance begins at point A on the phase diagram below and the pressure on the substance is reduced at constant temperature until point B is reached, what process occurs?
A. fusion B. vaporization C. condensation D. sublimation E. none of the above
24. What process occurs when a substance is at point C on the phase diagram below, and the pressure is decreased (under constant temperature) until the substance is at point E?
A. condensation B. vaporization C. sublimation D. melting E. freezing
25. In the phase diagram below, which phase is most dense?
A. solid B. liquid C. gas D. supercritical fluid E. none of the above
26. Which of the following phase transitions might be observed as the temperature of a pure substance is increased under constant pressure? 1. 2. 3.
solid ® liquid liquid ® gas gas ® solid
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 27. A permanent gas is one that may not be compressed to a liquid at 25C. Which of the following are permanent gases? gas O2 NH3 SO2
A. O2 only B. NH3 only C. SO2 only D. O2 and NH3 E. NH3 and SO2
b.p. (C) -183 -33.3 -10
crit. temp.(C) -119 31 158
28. All of the following statements concerning dispersion forces are correct EXCEPT A. the strength of dispersion forces depends on the number of electrons in a molecule. B. dispersion forces are the primary attractive forces in metal-metal bonding. C. dispersion forces involve the attraction between induced dipoles. D. dispersion forces are the primary attractive force in molecular solids consisting of nonpolar molecules. E. dispersion forces exist in all molecular solids.
29. The boiling points of some group 15 hydrides are tabulated below.
gas NH 3 PH 3 AsH 3
b.p. (C) -33.3 -88 -63
Which intermolecular force or bond is responsible for the high boiling point of ammonia relative to the other hydrides?
A. dispersion forces B. dipole forces C. hydrogen bonding D. covalent bonding E. ionic bonding 30. All of the following statements concerning hydrogen bonding are correct EXCEPT A. water is less dense in the solid phase than the liquid phase due to hydrogen bonding. B. hydrogen bonding occurs in molecules with N-H, O-H, and F-H bonds. C. hydrogen bonding in water results in lower than expected melting points. D. the unusually high boiling points of water, ammonia, and hydrogen fluoride are the result of hydrogen bonding. E. hydrogen bonding is an unusually strong dipole force.
31. Which one of the following molecules will exhibit dipole forces as a pure liquid or solid? A. CS2 B. C2H2 C. CCl4 D. Br2 E. PH3
32. Which intermolecular forces is/are present in solid SO3? 1. 2. 3.
dispersion dipole hydrogen bonding
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1 and 3 33. What is the dominant intermolecular force in HF(l)? A. dispersion forces B. hydrogen bonding C. ionic bonding D. dipole forces E. induced dipole forces
34. Which of the following pure substances will have hydrogen bonds? (Lone electron pairs have been omitted from these structures.)
A. acetone B. dimethyl ether C. methanol D. acetone and methanol E. dimethyl ether and methanol
35. Arrange NH3, CH4, and PH3 in order from lowest to highest boiling points. A. CH4 < PH3 < NH3 B. CH4 < NH3 < PH3 C. PH3 < NH3 < CH4 D. NH3 < CH4 < PH3 E. PH3 < CH4 < NH3
36. Which one of the following molecules has the lowest boiling point? A. CH4 B. CHCl3 C. CH2Cl2 D. CH3Cl E. CCl4
37. Which of the following nonpolar molecules has the highest boiling point? A. F2 B. C2H4 C. F2 D. O2 E. CS2
38. Arrange N2, O2, He, and Cl2 in order from lowest to highest melting point. A. O2 < N2 < Cl2 < He B. N2 < O2 < He < Cl2 C. He < Cl2 < N2 < O2 D. He < O2 < N2 < Cl2 E. He < N2 < O2 < Cl2
39. Arrange the noble gases in order from weakest to strongest interatomic forces. A. Kr < Ar < Ne < He B. Ne < He <Kr < Ar C. He < Ar < Ne < Kr D. He < Ne < Ar < Kr E. Ar < Kr < He < Ne
40. Which of the following substances will exhibit dipole forces? A. SO3 B. H2S C. CH4 D. SF6 E. N2
41. All of the following substances will exhibit dipole forces EXCEPT A. SO3. B. SF2. C. H2Se. D. NO2. E. NO.
42. In which of the following pure solids is it necessary to break covalent bonds in order to make a liquid or gas? A. SO2 B. NaCl C. H2SO4 D. I2 E. SiO2
43. Elemental phosphorus is a molecular solid consisting of P4 molecules. It melts at 44C. What is the principal force present in P4(s)? A. dipole forces B. hydrogen bonding C. dispersion forces D. metallic bonding E. ionic bonding
44. Arrange Cl2, ICl, and Br2 in order from lowest to highest melting point. A. Br2 < ICl < Cl2 B. Br2 < Cl2 < ICl C. Cl2 < ICl < Br2 D. Cl2 < Br2 < ICl E. ICl < Br2 < Cl2
45. All of the following statements are correct EXCEPT A. network covalent solids are usually good electrical conductors. B. network covalent solids often have high melting points. C. network covalent solids are insoluble in common solvents. D. ionic solids typically have high melting points. E. ionic solids are often hard and brittle.
46. Which of the following is NOT a network covalent solid? A. elemental silicon, Si(s) B. diamond, C(s) C. buckminster fullerene, C60(s) D. silicon dioxide, SiO2(s) E. aluminum oxide, Al2O3(s)
47. How many unit cells share an atom which is located at a corner (or lattice point) of a unit cell? A. 1 B. 2 C. 4 D. 6 E. 8
48. Which of the following statements concerning a metal crystallized in a face-centered cubic cell is/are CORRECT? 1. 2. 3.
One metal atom is located on each face of the unit cell, where it is shared equally between four unit cells. One metal atom is located at the center of the unit cell. A metal atom is located at each of the eight lattice points, where it is shared equally between eight unit cells.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2 and 3 49. What is a correct method for determining the number of atoms in a body-centered cubic cell? A. B. C. D. E.
50. What is the correct method for determining the number of atoms in a face-centered cubic cell? A. B. C. D. E.
51. Arrange the three common unit cells in order from least dense to most dense packing. A. simple cubic < body-centered cubic < face-centered cubic B. simple cubic < face-centered cubic < body-centered cubic C. body-centered cubic < face-centered cubic < simple cubic D. body-centered cubic < simple cubic < face-centered cubic E. face-centered cubic < body-centered cubic < simple cubic
52. Iron packs in a body-centered cubic structure. If an iron atom has a radius of 126 pm, what is the distance between lattice points at two opposite corners of the unit cell? (Note: a line drawn between the points would go through the center of the unit cell) A. B. 2(126 pm) = 252 pm
C. D. 4(126 pm) = 504 pm E.
53. What is the distance, in atomic radii, along any edge of a face-centered unit cell? A. B. 2 ´ r C. 4 ´ r D. E. r
54. For a simple cubic unit cell, what percentage of the space in the cell is occupied by the atoms at the corners of the cell? A. 47.6% B. 52.4% C. 57.4% D. 62.3% E. 71.2%
55. Nickel has a face-centered cubic cell, and its density is 8.90 g/cm3. What is the radius (in pm) of a nickel atom? (The molar mass of nickel is 58.69 g/mol) A. 62.3 pm B. 88.1 pm C. 125 pm D. 249 pm E. 535 pm
56. Rhodium (atomic mass 102.9 g/mol) crystallizes in a face-centered cubic unit cell. In addition, rhodium has an atomic radius of 135 pm. What is the density (in g/cm3) of rhodium? A. 1.53 g/cm3 B. 6.14 g/cm3 C. 17.4 g/cm3 D. 12.3 g/cm3 E. 27.8 g/cm3
57. Potassium crystallizes in a body-centered cubic unit cell. If the length of an edge of the unit cell is 524 pm, what is the atomic radius (in pm) of a potassium atom? A. 151 pm B. 185 pm C. 227 pm D. 262 pm E. 371 pm
58. Niobium crystallizes in a body-centered cubic unit cell. If the radius of a niobium atom is 0.145 nm, what is the length of an edge of the unit cell? A. 0.251 nm B. 0.290 nm C. 0.335 nm D. 0.410 nm E. 0.502 nm
59. Dorothy Crowfoot Hodgkin was famous for what? A. the discovery of x-rays B. the invention of x-ray diffraction C. the discovery of supercritical fluids D. the determination of the structure of insulin E. all of the above
60. What type of solid is malleable? A. metallic B. ionic C. molecular D. crystalline E. covalent
61. What type of solid is generally soluble in nonpolar solvents? A. metallic B. ionic C. molecular D. crystalline E. covalent
62. What type of solid is conductive when melted but not as a solid? A. metallic B. ionic C. molecular D. crystalline E. covalent
Chapter 9--Liquids and Solids Key
1. Equilibrium has been established between a liquid and its vapor when A. the masses of liquid and vapor are equal. B. all of the liquid has evaporated. C. when the temperatures of the liquid and the vapor are the same. D. evaporation ceases and the concentrations of liquid and vapor remain constant. E. the rate of condensation equals the rate of evaporation.
2. Which of the following statements is/are CORRECT? The vapor pressure of a liquid depends on 1. 2. 3.
the temperature of the liquid. the surface area of the liquid. the volume of the liquid.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 3. Methyl alcohol, CH3OH, has a vapor pressure of 203 mm Hg at 35C. If 5.00 g CH3OH is sealed in a 10.0 L flask, what mass will remain in the liquid phase when equilibrium is established at 35C? Assume any liquid remaining in the flask has a negligible volume. (760 mm Hg = 1 atm, R = 0.0821 Latm/molK) A. 0.00 g B. 1.62 g C. 2.08 g D. 2.27 g E. 4.69 g
4. At 75.0 C, water has an equilibrium vapor pressure of 289.1 mm Hg. If 4.22 g H2O is sealed in an evacuated 5.00 L flask and heated to 75.0 C, what mass of H2O will be found in the gas phase when liquid-vapor equilibrium is established? Assume any liquid remaining in the flask has a negligible volume. (760 mm Hg = 1 atm, R = 0.0821 Latm/molK) A. 0.240 g B. 1.20 g C. 2.64 g D. 3.02 g E. 4.22 g
5. Water has an equilibrium vapor pressure of 23.8 mm Hg at 25C. What mass of water vapor is present, at equilibrium, in a room with dimensions of 7.0 m ´ 8.0 m ´ 2.7 m? (760 mm Hg = 1 atm, R = 0.0821 Latm/molK) A. 0.0050 kg B. 0.029 kg C. 0.28 kg D. 3.5 kg E. 5.0 kg
6. Which of the following statements is true for an ideal gas? A. A plot of ln P vsersus 1/T (in Kelvin) yields a straight line with a slope equal to -DHvap. B. A plot of ln P versus T (in Kelvin) yields a straight line with a slope equal to DHvap. C. A plot of ln P versus 1/T (in Kelvin) yields a straight line with a slope equal to -DHvap/R. D. A plot of P versus 1/T (in Kelvin) yields a straight line with a slope equal to 1/-DHvap. E. A plot of P versus T (in Kelvin) yields a straight line with a slope equal to -DHvap/RT.
7. Which of the following equations is a correct form of the Clausius-Clapeyron equation?
A.
B.
C.
D.
E.
8. Sulfur dioxide has a vapor pressure of 462.7 mm Hg at –21.0C and a vapor pressure of 140.5 mm Hg at –44.0C. What is the molar heat of vaporization of sulfur dioxide? (R = 8.31 J/Kmol) A. 0.398 kJ/mol B. 6.33 kJ/mol C. 14.0 kJ/mol D. 24.9 kJ/mol E. 39.8 kJ/mol
9. Ethanol has a molar heat of vaporization of 42.3 kJ/mol. The compound has a vapor pressure of 1.00 atm at 78.3C. At what temperature is the vapor pressure equal to 0.800 atm? (R = 8.31 J/Kmol) A. –83.8C B. –24.4C C. 62.6C D. 73.0C E. 78.0C
10. Carbon tetrachloride, an organic solvent, has a molar heat of vaporization of 29.82 kJ/mol. If CCl4 has a normal boiling point of 3.50 ´ 102 K, what is its vapor pressure at 273 K? (R = 8.31 J/molK) A. 19.3 mm Hg B. 42.2 mm Hg C. 59.3 mm Hg D. 108 mm Hg E. 359 mm Hg
11. Mount Everest rises to a height of 29,035 ft (8.850 ´ 103 m) above sea level. At this height, the atmospheric pressure is 230 mm Hg. At what temperature does water boil at the summit of Mount Everest? (DHvap for H2O = 40.7 kJ/mole, R = 8.31 J/Kmol) A. -17C B. 31C C. 48C D. 57C E. 69C
12. Sulfur dioxide has an enthalpy of vaporization of 24.9 kJ/mol. At 205 K, SO2 has a vapor pressure of 30.3 mm Hg. What is the normal boiling point temperature of SO2? (R = 8.31 J/Kmol) A. 263 K B. 278 K C. 291 K D. 308 K E. 345 K
13. Freon-113, C2Cl3F3, has an enthalpy of vaporization of 27.04 kJ/mol and a normal boiling point of 48C. At what temperature is the vapor pressure of Freon-113 equal to 76.0 mm Hg? (R = 8.31 J/Kmol) A. -53.1C B. -23.8C C. -11.4C D. 1.44C E. 7.12C
14. Xenon has a molar heat of vaporization of 12.6 kJ/mol and a vapor pressure of 1.00 atm at –108.0C. What is the vapor pressure of xenon at -148.0C? (R = 8.31 J/Kmol) A. 0.053 atm B. 0.73 atm C. 0.93 atm D. 0.99 atm E. 19 atm
15. Naphthalene, a substance present in some mothballs, has a molar heat of vaporization of 49.4 kJ/mol. If the vapor pressure of naphthalene is 0.300 mm Hg at 298 K, what is the temperature at which the vapor pressure is 7.60 ´ 102 mm Hg? (R = 8.31 J/Kmol) A. 424 K B. 491 K C. 501 K D. 521 K E. 611 K
16. The normal boiling point of a liquid is A. 373 K. B. the temperature at which a liquid's vapor pressure equals 1 atm. C. the pressure at which the liquid boils at 373 K. D. dependent upon the volume of the liquid. E. dependent upon the surface area of the liquid..
17. All of the following statements are incorrect EXCEPT A. gas, liquid and solid phases of a substance exist simultaneously at equilibrium at the critical pressure and temperature. B. above the critical pressure, only the solid phase of a pure substance can exist. C. the liquid phase of a pure substance cannot exist above its critical temperature. D. the critical temperature and pressure cannot be produced simultaneously. E. above the critical temperature and pressure, only the liquid phase of a substance exists.
18. Which of the following statements are correct? 1. 2. 3.
All three phases (gas, liquid, and solid) exist at equilibrium at the triple point. Above its critical temperature and pressure, a substance becomes a supercritical fluid. A liquid boils when its vapor pressure equals the pressure above its surface.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 19. The phase diagram for CO2 has a triple point at –56.6 C and 5.19 atm, and a critical point at 31.0 C and 73 atm. The solid and gas phases are in equilibrium at –78.7 C and 1.00 atm. Which of the following statements regarding CO2 is/are CORRECT? 1. 2. 3.
Sublimation occurs if the temperature of the solid phase is increased from –79.0 C to 0.0 C at a constant pressure of 2.5 atm. CO2 is a supercritical fluid at 55 C and 75 atm. At pressures greater than its critical pressure (73 atm), CO2 will not exist as a solid at any temperature.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 20. On the phase diagram below, which point corresponds to conditions where only the solid phase exists?
A. A B. B C. C D. D E. G
21. On the phase diagram below, which point corresponds to conditions where both liquid and gas phases exist?
A. A B. B C. C D. E E. F
22. On the phase diagram below, which point corresponds to conditions where solid, liquid, and gas phases all exist?
A. B B. C C. D D. E E. G
23. If a pure substance begins at point A on the phase diagram below and the pressure on the substance is reduced at constant temperature until point B is reached, what process occurs?
A. fusion B. vaporization C. condensation D. sublimation E. none of the above
24. What process occurs when a substance is at point C on the phase diagram below, and the pressure is decreased (under constant temperature) until the substance is at point E?
A. condensation B. vaporization C. sublimation D. melting E. freezing
25. In the phase diagram below, which phase is most dense?
A. solid B. liquid C. gas D. supercritical fluid E. none of the above
26. Which of the following phase transitions might be observed as the temperature of a pure substance is increased under constant pressure? 1. 2. 3.
solid ® liquid liquid ® gas gas ® solid
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 27. A permanent gas is one that may not be compressed to a liquid at 25C. Which of the following are permanent gases? gas O2 NH3 SO2
A. O2 only B. NH3 only C. SO2 only D. O2 and NH3 E. NH3 and SO2
b.p. (C) -183 -33.3 -10
crit. temp.(C) -119 31 158
28. All of the following statements concerning dispersion forces are correct EXCEPT A. the strength of dispersion forces depends on the number of electrons in a molecule. B. dispersion forces are the primary attractive forces in metal-metal bonding. C. dispersion forces involve the attraction between induced dipoles. D. dispersion forces are the primary attractive force in molecular solids consisting of nonpolar molecules. E. dispersion forces exist in all molecular solids.
29. The boiling points of some group 15 hydrides are tabulated below.
gas NH 3 PH 3 AsH 3
b.p. (C) -33.3 -88 -63
Which intermolecular force or bond is responsible for the high boiling point of ammonia relative to the other hydrides?
A. dispersion forces B. dipole forces C. hydrogen bonding D. covalent bonding E. ionic bonding 30. All of the following statements concerning hydrogen bonding are correct EXCEPT A. water is less dense in the solid phase than the liquid phase due to hydrogen bonding. B. hydrogen bonding occurs in molecules with N-H, O-H, and F-H bonds. C. hydrogen bonding in water results in lower than expected melting points. D. the unusually high boiling points of water, ammonia, and hydrogen fluoride are the result of hydrogen bonding. E. hydrogen bonding is an unusually strong dipole force.
31. Which one of the following molecules will exhibit dipole forces as a pure liquid or solid? A. CS2 B. C2H2 C. CCl4 D. Br2 E. PH3
32. Which intermolecular forces is/are present in solid SO3? 1. 2. 3.
dispersion dipole hydrogen bonding
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1 and 3 33. What is the dominant intermolecular force in HF(l)? A. dispersion forces B. hydrogen bonding C. ionic bonding D. dipole forces E. induced dipole forces
34. Which of the following pure substances will have hydrogen bonds? (Lone electron pairs have been omitted from these structures.)
A. acetone B. dimethyl ether C. methanol D. acetone and methanol E. dimethyl ether and methanol
35. Arrange NH3, CH4, and PH3 in order from lowest to highest boiling points. A. CH4 < PH3 < NH3 B. CH4 < NH3 < PH3 C. PH3 < NH3 < CH4 D. NH3 < CH4 < PH3 E. PH3 < CH4 < NH3
36. Which one of the following molecules has the lowest boiling point? A. CH4 B. CHCl3 C. CH2Cl2 D. CH3Cl E. CCl4
37. Which of the following nonpolar molecules has the highest boiling point? A. F2 B. C2H4 C. F2 D. O2 E. CS2
38. Arrange N2, O2, He, and Cl2 in order from lowest to highest melting point. A. O2 < N2 < Cl2 < He B. N2 < O2 < He < Cl2 C. He < Cl2 < N2 < O2 D. He < O2 < N2 < Cl2 E. He < N2 < O2 < Cl2
39. Arrange the noble gases in order from weakest to strongest interatomic forces. A. Kr < Ar < Ne < He B. Ne < He <Kr < Ar C. He < Ar < Ne < Kr D. He < Ne < Ar < Kr E. Ar < Kr < He < Ne
40. Which of the following substances will exhibit dipole forces? A. SO3 B. H2S C. CH4 D. SF6 E. N2
41. All of the following substances will exhibit dipole forces EXCEPT A. SO3. B. SF2. C. H2Se. D. NO2. E. NO.
42. In which of the following pure solids is it necessary to break covalent bonds in order to make a liquid or gas? A. SO2 B. NaCl C. H2SO4 D. I2 E. SiO2
43. Elemental phosphorus is a molecular solid consisting of P4 molecules. It melts at 44C. What is the principal force present in P4(s)? A. dipole forces B. hydrogen bonding C. dispersion forces D. metallic bonding E. ionic bonding
44. Arrange Cl2, ICl, and Br2 in order from lowest to highest melting point. A. Br2 < ICl < Cl2 B. Br2 < Cl2 < ICl C. Cl2 < ICl < Br2 D. Cl2 < Br2 < ICl E. ICl < Br2 < Cl2
45. All of the following statements are correct EXCEPT A. network covalent solids are usually good electrical conductors. B. network covalent solids often have high melting points. C. network covalent solids are insoluble in common solvents. D. ionic solids typically have high melting points. E. ionic solids are often hard and brittle.
46. Which of the following is NOT a network covalent solid? A. elemental silicon, Si(s) B. diamond, C(s) C. buckminster fullerene, C60(s) D. silicon dioxide, SiO2(s) E. aluminum oxide, Al2O3(s)
47. How many unit cells share an atom which is located at a corner (or lattice point) of a unit cell? A. 1 B. 2 C. 4 D. 6 E. 8
48. Which of the following statements concerning a metal crystallized in a face-centered cubic cell is/are CORRECT? 1. 2. 3.
One metal atom is located on each face of the unit cell, where it is shared equally between four unit cells. One metal atom is located at the center of the unit cell. A metal atom is located at each of the eight lattice points, where it is shared equally between eight unit cells.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2 and 3 49. What is a correct method for determining the number of atoms in a body-centered cubic cell? A. B. C. D. E.
50. What is the correct method for determining the number of atoms in a face-centered cubic cell? A. B. C. D. E.
51. Arrange the three common unit cells in order from least dense to most dense packing. A. simple cubic < body-centered cubic < face-centered cubic B. simple cubic < face-centered cubic < body-centered cubic C. body-centered cubic < face-centered cubic < simple cubic D. body-centered cubic < simple cubic < face-centered cubic E. face-centered cubic < body-centered cubic < simple cubic
52. Iron packs in a body-centered cubic structure. If an iron atom has a radius of 126 pm, what is the distance between lattice points at two opposite corners of the unit cell? (Note: a line drawn between the points would go through the center of the unit cell) A. B. 2(126 pm) = 252 pm
C. D. 4(126 pm) = 504 pm E.
53. What is the distance, in atomic radii, along any edge of a face-centered unit cell? A. B. 2 ´ r C. 4 ´ r D. E. r
54. For a simple cubic unit cell, what percentage of the space in the cell is occupied by the atoms at the corners of the cell? A. 47.6% B. 52.4% C. 57.4% D. 62.3% E. 71.2%
55. Nickel has a face-centered cubic cell, and its density is 8.90 g/cm3. What is the radius (in pm) of a nickel atom? (The molar mass of nickel is 58.69 g/mol) A. 62.3 pm B. 88.1 pm C. 125 pm D. 249 pm E. 535 pm
56. Rhodium (atomic mass 102.9 g/mol) crystallizes in a face-centered cubic unit cell. In addition, rhodium has an atomic radius of 135 pm. What is the density (in g/cm3) of rhodium? A. 1.53 g/cm3 B. 6.14 g/cm3 C. 17.4 g/cm3 D. 12.3 g/cm3 E. 27.8 g/cm3
57. Potassium crystallizes in a body-centered cubic unit cell. If the length of an edge of the unit cell is 524 pm, what is the atomic radius (in pm) of a potassium atom? A. 151 pm B. 185 pm C. 227 pm D. 262 pm E. 371 pm
58. Niobium crystallizes in a body-centered cubic unit cell. If the radius of a niobium atom is 0.145 nm, what is the length of an edge of the unit cell? A. 0.251 nm B. 0.290 nm C. 0.335 nm D. 0.410 nm E. 0.502 nm
59. Dorothy Crowfoot Hodgkin was famous for what? A. the discovery of x-rays B. the invention of x-ray diffraction C. the discovery of supercritical fluids D. the determination of the structure of insulin E. all of the above
60. What type of solid is malleable? A. metallic B. ionic C. molecular D. crystalline E. covalent
61. What type of solid is generally soluble in nonpolar solvents? A. metallic B. ionic C. molecular D. crystalline E. covalent
62. What type of solid is conductive when melted but not as a solid? A. metallic B. ionic C. molecular D. crystalline E. covalent
Chapter 10--Solutions 1. What is the definition of molarity? A. mass of solute per liter of solvent B. mass of solute per kg of solvent C. moles of solute per kg of solvent D. moles of solute in one liter of solvent E. moles of solute per liter of solution
2. To prepare 0.250 L of 0.100 M aqueous NaCl (58.4 g/mol), one may A. dissolve 0.100 g of NaCl in 250 mL of water. B. dissolve 1.46 g of NaCl in 250 mL of water. C. dissolve 0.100 g of NaCl in enough water to make 0.250 kg of solution. D. dissolve 1.46 g of NaCl in enough water to make 0.250 L of solution. E. dissolve 0.100 g NaCl in 0.250 kg of water.
3. If 25.00 mL of 2.00 M NaCl is transferred by pipet into a volumetric flask and diluted to 5.00 L, what is the molarity of the diluted NaCl? A. 0.0100 M B. 0.0160 M C. 0.0625 M D. 0.400 M E. 16.0 M
4. What volume of 6.0 M HNO3 is required to prepare 250 mL of 0.40 M HNO3? A. 9.7 mL B. 17 mL C. 27 mL D. 38 mL E. 270 L
5. If 26.5 g of methanol (CH3OH) is added to 735 g of water, what is the molality of the methanol? A. 0.0348 m B. 2.03 m C. 1.13 m D. 3.61 m E. 36.1 m
6. What mass of Cu(NO3)2 (187.6 g/mol) is present in 25.0 g of 1.00 m Cu(NO3)2(aq)? A. 3.95 g B. 4.69 g C. 13.8 g D. 25.0 g E. 63.5 g
7. To prepare a solution that is 15.0% aqueous KCl by mass, one should A. dissolve 15.0 g KCl in 85.0 g H2O. B. dissolve 15.0 g KCl in 1.00 ´ 102 g H2O. C. dissolve 15.0 g KCl in 0.850 mol H2O. D. dissolve 0.150 mol KCl in 0.850 mol H2O. E. dissolve 0.150 mol KCl in 1.00 mol H2O.
8. What mass of HCl is required to prepare 1.00 kg of 5.5% by mass aqueous HCl? A. 0.018 g B. 5.5 g C. 18 g D. 55 g E. 550 g
9. What is the mole fraction of water in a solution that is 33.3% by mass ethylene glycol? The molar mass of ethylene glycol, HOCH2CH2OH, is 62.07 g/mol. A. 0.127 B. 0.290 C. 0.368 D. 0.667 E. 0.873
10. The mole fraction of calcium chloride in an aqueous solution is 0.0724. What is the percent mass of CaCl2 in the solution? A. 4.46% B. 7.24% C. 32.5% D. 36.2% E. 50.0%
11. Concentrated nitric acid is 70.4% HNO3 by mass. What is the mole fraction of nitric acid? A. 0.0112 B. 0.0620 C. 0.171 D. 0.377 E. 0.405
12. Concentrated sodium hydroxide is 19.4 M and has a density of 1.54 g/mL. What is the molality of concentrated NaOH? A. 12.6 m B. 19.8 m C. 25.4 m D. 29.9 m E. 50.4 m
13. Concentrated sulfuric acid is 18.0 M and has a density of 1.84 g/mL. Calculate the percent mass of sulfuric acid in concentrated H2SO4. A. 17.7% B. 32.5% C. 78.2% D. 96.0% E. 99.4%
14. A bottle of phosphoric acid is labeled "85.0% H3PO4 by mass; density = 1.689 g/cm3." Calculate the molarity of phosphoric acid in the solution. A. 16.7 M B. 14.6 M C. 12.2 M D. 8.67 M E. 0.146 M
15. Concentrated phosphoric acid is 85.0% by mass H3PO4. If the molarity of concentrated H3PO4 is 14.5 M, what is the density? A. 0.60 g/mL B. 1.67 g/mL C. 1.87 g/mL D. 1.95 g/mL E. 2.07 g/mL
16. Pure acetic acid, often called glacial acetic acid, is a liquid with a density of 1.049 g/mL. Which calculation correctly shows how to determine the volume of glacial acetic acid necessary to prepare 250 mL of 0.400 M CH3CO2H(aq)?
A.
B.
C.
D.
E.
17. Pure acetic acid, CH3CO2H(l), has a density of 1.049 g/mL. To prepare 1.00 L of 6.00 M CH3CO2H(aq), one may A. dilute 175 g of acetic acid to a volume of 1.00 L. B. dilute 343 mL of acetic acid to a volume of 1.00 L. C. dilute 360 mL of acetic acid to a volume of 1.00 L. D. dilute 382 mL of acetic acid to a volume of 1.00 L. E. dilute 1049 g of acetic acid to a volume of 1.00 L.
18. A 15 meter by 12 meter pool of water has a depth of 2.2 meters. What mass of silver ion is present in the reservoir if the concentration of silver ion is 0.14 ppm? (1 m3 = 1000 L; assume the density of the solution is 1.00 g/mL) A. 5.5 ´ 10–4 g B. 5.5 ´ 10–2 g C. 0.55 g D. 5.5 g E. 55 g
19. If sea water contains 15 ppm gold, how many kilograms of sea water must be processed to remove 1.00 g of gold? A. 67 kg B. 97 kg C. 150 kg D. 6.7 ´ 103 kg E. 15 ´ 104 kg
20. Silver chloride is a relatively insoluble salt. Only 1.92 mg of AgCl will dissolve per liter of water at 25C. How many parts per million of Ag+ can be present in water at 25C? Assume the density of the solution equals the density of water, 1.00 g/mL. A. 0.52 ppm B. 0.93 ppm C. 1.45 ppm D. 9.30 ppm E. 52.0 ppm
21. Silver chloride is a relatively insoluble salt. Only 1.92 mg of AgCl will dissolve per liter of water at 25C. What concentration of Ag+, in molarity units, can be present in water at 25C? A. 3.35 ´ 10-6 M B. 6.70 ´ 10-6 M C. 1.34 ´ 10-5 M D. 9.60 ´ 10-4 M E. 1.92 ´ 10-3 M
22. What concentration of silver nitrate (in ppm) is present in 7.1 ´ 10–7 M AgNO3(aq)? For very dilute aqueous solutions, you can assume the solution's density is 1.0 g/mL. The molar mass of AgNO3 is 169.9 g/mol. A. 0.0071 ppm B. 0.12 ppm C. 0.71 ppm D. 1.7 ppm E. 8.3 ppm
23. A gas mixture has mole fractions of 0.24 oxygen and 0.76 nitrogen. If the total pressure of the gases is 1.44 atm at 325 K, what is the concentration, in molarity, of oxygen? (R = 0.0821 Latm/molK) A. 0.013 M B. 0.017M C. 0.041 M D. 0.54 M E. 0.069 M
24. If 77.5 g of ethylene glycol (HOCH2CH2OH) is added to 422.5 g of water, what is the mole fraction of ethylene glycol? A. 0.00296 B. 0.0506 C. 0.183 D. 0.949 E. 2.96
25. Arrange the molecules below in order of increasing solubility in water.
A. dimethyl ether < methane < methanol B. dimethyl ether < methanol < methane C. methane < methanol< dimethyl ether D. methanol < dimethyl ether < methane E. methane < dimethyl ether < methanol
26. Which of the following liquids will be miscible with water in any proportions: ethanol (CH3CH2OH), carbon tetrachloride (CCl4), hexane (C6H14), and/or formic acid (HCO2H)? A. ethanol and carbon tetrachloride B. carbon tetrachloride and hexane C. ethanol and formic acid D. ethanol, carbon tetrachloride, and benzene E. carbon tetrachloride, and formic acid
27. Arrange the molecules below in order of increasing solubility in water.
A. methane < methanol < chloromethane B. methane < chloromethane < methanol C. chloromethane < methanol < methane D. chloromethane < methane < methanol E. methanol < methane < chloromethane
28. All of the following statements are correct EXCEPT A. the solubility of a gas in water decreases as the water temperature increases. B. dissolving a solid in water is usually an exothermic process. C. when an equilibrium is established between molecules in a solid and a solution, the solution is said to be saturated. D. if a precipitate forms when a solution is cooled, the solution is supersaturated. E. network covalent solids are usually insoluble in water.
29. A substance that dissolves in water and conducts electricity when present in an aqueous solution is A. metallic. B. ionic. C. molecular. D. network covalent. E. none of the above.
30. Which of the following statements concerning the solubility of a gas in a liquid are true? 1. 2. 3.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1 and 3
Solubility decreases with increasing temperature. Solubility increases as the pressure of the gas over the liquid increases. Solubility is dependent on the surface area of the the liquid.
31. Henry's law states that gas solubility is A. directly proportional to temperature of the solution. B. directly proportional to the molar mass of the gas. C. inversely proportional to the combined pressure of all gases over the solution. D. inversely proportional to the pressure of the gas over the solution. E. directly proportional to the pressure of the gas over the solution.
32. The Henry's law constant for O2 in water at 25 C is 1.26 ´ 10–3 M/atm. What partial pressure of O2 is necessary to achieve an equilibrium concentration of 1.5 ´ 10–3 M O2? A. 1.9 ´ 10–6 atm B. 0.24 atm C. 0.84 atm D. 1.2 atm E. 5.3 ´ 105 atm
33. The Henry's law constant for the solubility of nitrogen in water is 6.4 ´ 10-4 M/atm at 25C. At 0.75 atm of N2, what mass of N2(g) dissolves in 1.0 L of water at 25C? A. 4.8 ´ 10-4 g B. 8.5 ´ 10-4 g C. 4.5 ´ 10-3 g D. 1.3 ´ 10-2 g E. 2.4 ´ 10-2 g
34. In 1.00 atm of pure oxygen, the solubility of O2(g) in water is 1.26 ´ 10-3 M at 25.0C. The mole fraction of oxygen in air is 0.210. If the atmospheric pressure is 0.979 atm, what is the solubility of oxygen in air at 25.0C? A. 2.59 ´ 10-4 M B. 2.65 ´ 10-3 M C. 5.87 ´ 10-3 M D. 6.00 ´ 10-3 M E. 3.86 ´ 103 M
35. At 1.00 atm, 1.64 ´ 10-3 g H2(g) will dissolve in 1.0 L of water. What pressure of gas is necessary to obtain a concentration of 1.0 ´ 10-3 M H2(g)? A. 6.1 ´ 10-4 atm B. 4.4 ´ 10-2 atm C. 1.2 atm D. 2.2 ´ 103 atm E. 1.6 ´ 103 atm
36. All of the following are colligative properties EXCEPT A. gas solubility as a function of partial pressure above a solution (Henry's law). B. osmotic pressure. C. vapor pressure lowering. D. boiling point elevation. E. freezing point depression.
37. An aqueous solution contains 11.5 g of NaCl in 250.2 g of water. Calculate the vapor pressure of this solution at 25.0C. The vapor pressure of pure water is 23.8 mm Hg at 25.0C. A. 3.38 mm Hg B. 21.8 mm Hg C. 23.5 mm Hg D. 24.1 mm Hg E. 31.3 mm Hg
38. What is the equilibrium partial pressure of water vapor above a mixture of 44.0 g H2O and 56.0 g HOCH2CH2OH at 35C. The partial pressure of pure water at 35.0C is 42.2 mm Hg. Assume ideal behavior for the solution. A. 0.730 Hg B. 18.7 mm Hg C. 23.6 mm Hg D. 30.8 mm Hg E. 58.8 mm Hg
39. What mass of ethylene glycol, when mixed with 225 g H2O, will reduce the equilibrium vapor pressure of H2O from 1.00 atm to 0.800 atm at 100C? The molar masses of water and ethylene glycol are 18.02 g/mol and 62.07 g/mol, respectively. Assume ideal behavior for the solution. A. 15.6 g B. 49.9 g C. 194 g D. 969 g E. 3.10 ´ 103 g
40. Which of the following aqueous solutions will freeze at the lowest temperature? A. 0.10 m KCl B. 0.20 m C6H12O6 (glucose) C. 0.050 m AlCl3 D. 0.15 m SrBr2 E. All of the above freeze at the same temperature.
41. What concentration unit is necessary for the calculation of boiling point elevation? A. molarity B. molality C. mass fraction D. mole fraction E. parts per million
42. What concentration unit is necessary for the calculation of vapor pressure lowering? A. molarity B. molality C. mass fraction D. mole fraction E. density
43. The molal boiling point constant for water is 0.52C/m. At what temperature will a mixture of 45.0 g of NaCl and 0.500 kg of water boil? A. 98.4C B. 99.2C C. 100.0C D. 100.8C E. 101.6C
44. What is the boiling point of a solution containing 0.80 g caffeine, C8H10N4O2, dissolved in 13.20 g benzene? The boiling point of pure benzene is 80.1C and the molal boiling point constant, Kb, is 2.53C/m. A. 79.8C B. 80.4C C. 80.9C D. 85.2C E. 88.2C
45. The molal freezing point constant for water is –1.86C/m. At what temperature will a solution containing 8.27 g CaCl2 and 45.0 g H2O begin to freeze? Assume that no ion-pairing occurs between Ca2+ and Cl–. A. –9.24C B. –4.62C C. –0.804C D. –0.749C E. +4.62C
46. What is the molar mass of a nonpolar molecular compound if 3.42 grams dissolved in 41.8 grams benzene begins to freeze at 1.17C? The freezing point of pure benzene is 5.50C and the molal freezing point constant, Kf, is –5.12C/m. A. 2.89 g/mol B. 69.2 g/mol C. 96.7 g/mol D. 126 g/mol E. 358 g/mol
47. Equal masses of water and ethylene glycol (C2H6O2) are mixed. At what temperature will the mixture freeze? The molal freezing point constant for water is -1.86C/m. A. -115C B. -93.0C C. -42.0C D. -30.0C E. -0.93C
48. Which of the following statements concerning osmosis is/are CORRECT? 1. 2. 3.
Osmosis involves the movement of a solvent through a semipermeable membrane. Solvents move from regions of high solute concentration to regions of lower solute concentration. Osmotic pressure is a colligative property.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 49. At 25C, what is the osmotic pressure of a homogeneous solution consisting of 18.0 g urea (CON2H4) diluted with water to 3.00 L? (R = 0.0821 Latm/molK) A. 0.205 atm B. 2.44 atm C. 7.33 atm D. 12.3 atm E. 14.7 atm
50. An aqueous solution is composed of 7.50 g NaCl (MM = 58.44 g/mol) diluted to 0.100 L. Calculate the osmotic pressure of the solution at 298 K. (R = 0.0821 Latm/molK) A. 5.83 atm B. 9.22 atm C. 18.3 atm D. 31.4 atm E. 62.8 atm
51. A solution is prepared by dissolving 5.00 g of an unknown molecular solid in water to make 1.00 L of solution. The osmotic pressure of the solution is 1.61 atm at 25C. What is the molar mass of the solute? (R = 0.0821 Latm/molK) A. 6.37 g/mol B. 58.44 g/mol C. 76.0 g/mol D. 102 g/mol E. 180.2 g/mol
52. A solution is prepared by dissolving 4.78 g of an unknown nonelectrolyte in enough water to make 375 mL of solution. The osmotic pressure of the solution is 1.33 atm at 27C. What is the molar mass of the solute? (R = 0.0821 Latm/molK) A. 0.0203 g/mol B. 21.2 g/mol C. 49.4 g/mol D. 96.8 g/mol E. 236 g/mol
53. A solution is prepared by dissolving 4.21 g of a nonelectrolyte in 50.0 g of water. If the boiling point increases by 0.203C, what is the molar mass of the solute? The boiling point elevation constant for water is 0.512C/m. A. 33.4 g/mol B. 111 g/mol C. 172 g/mol D. 212 g/mol E. 810 g/mol
54. Which of the following electrolytes is likely to have a van't Hoff factor equal to 3? A. CaI2 B. Na3PO4 C. KCl D. answers a and b E. answers a, b, and c
55. A 0.230 m solution of an unknown electrolyte depresses the freezing point of water by 0.821C. What is the Van't Hoff factor for this electrolyte? The freezing point depression constant for water is 1.86C/m. A. 0.521 B. 1.92 C. 2.00 D. 2.30 E. 4.41
56. Maple syrup is made from the sap of the maple tree. When the sap is tapped from the maple tree it is 2.0% by mass sucrose. Maple syrup is 66% by mass sucrose. What mass of sap would be required to make 0.5 kg of maple syrup? A. 0.50 kg B. 2.0 kg C. 66 kg D. 15 kg E. 17 kg
Chapter 10--Solutions Key
1. What is the definition of molarity? A. mass of solute per liter of solvent B. mass of solute per kg of solvent C. moles of solute per kg of solvent D. moles of solute in one liter of solvent E. moles of solute per liter of solution
2. To prepare 0.250 L of 0.100 M aqueous NaCl (58.4 g/mol), one may A. dissolve 0.100 g of NaCl in 250 mL of water. B. dissolve 1.46 g of NaCl in 250 mL of water. C. dissolve 0.100 g of NaCl in enough water to make 0.250 kg of solution. D. dissolve 1.46 g of NaCl in enough water to make 0.250 L of solution. E. dissolve 0.100 g NaCl in 0.250 kg of water.
3. If 25.00 mL of 2.00 M NaCl is transferred by pipet into a volumetric flask and diluted to 5.00 L, what is the molarity of the diluted NaCl? A. 0.0100 M B. 0.0160 M C. 0.0625 M D. 0.400 M E. 16.0 M
4. What volume of 6.0 M HNO3 is required to prepare 250 mL of 0.40 M HNO3? A. 9.7 mL B. 17 mL C. 27 mL D. 38 mL E. 270 L
5. If 26.5 g of methanol (CH3OH) is added to 735 g of water, what is the molality of the methanol? A. 0.0348 m B. 2.03 m C. 1.13 m D. 3.61 m E. 36.1 m
6. What mass of Cu(NO3)2 (187.6 g/mol) is present in 25.0 g of 1.00 m Cu(NO3)2(aq)? A. 3.95 g B. 4.69 g C. 13.8 g D. 25.0 g E. 63.5 g
7. To prepare a solution that is 15.0% aqueous KCl by mass, one should A. dissolve 15.0 g KCl in 85.0 g H2O. B. dissolve 15.0 g KCl in 1.00 ´ 102 g H2O. C. dissolve 15.0 g KCl in 0.850 mol H2O. D. dissolve 0.150 mol KCl in 0.850 mol H2O. E. dissolve 0.150 mol KCl in 1.00 mol H2O.
8. What mass of HCl is required to prepare 1.00 kg of 5.5% by mass aqueous HCl? A. 0.018 g B. 5.5 g C. 18 g D. 55 g E. 550 g
9. What is the mole fraction of water in a solution that is 33.3% by mass ethylene glycol? The molar mass of ethylene glycol, HOCH2CH2OH, is 62.07 g/mol. A. 0.127 B. 0.290 C. 0.368 D. 0.667 E. 0.873
10. The mole fraction of calcium chloride in an aqueous solution is 0.0724. What is the percent mass of CaCl2 in the solution? A. 4.46% B. 7.24% C. 32.5% D. 36.2% E. 50.0%
11. Concentrated nitric acid is 70.4% HNO3 by mass. What is the mole fraction of nitric acid? A. 0.0112 B. 0.0620 C. 0.171 D. 0.377 E. 0.405
12. Concentrated sodium hydroxide is 19.4 M and has a density of 1.54 g/mL. What is the molality of concentrated NaOH? A. 12.6 m B. 19.8 m C. 25.4 m D. 29.9 m E. 50.4 m
13. Concentrated sulfuric acid is 18.0 M and has a density of 1.84 g/mL. Calculate the percent mass of sulfuric acid in concentrated H2SO4. A. 17.7% B. 32.5% C. 78.2% D. 96.0% E. 99.4%
14. A bottle of phosphoric acid is labeled "85.0% H3PO4 by mass; density = 1.689 g/cm3." Calculate the molarity of phosphoric acid in the solution. A. 16.7 M B. 14.6 M C. 12.2 M D. 8.67 M E. 0.146 M
15. Concentrated phosphoric acid is 85.0% by mass H3PO4. If the molarity of concentrated H3PO4 is 14.5 M, what is the density? A. 0.60 g/mL B. 1.67 g/mL C. 1.87 g/mL D. 1.95 g/mL E. 2.07 g/mL
16. Pure acetic acid, often called glacial acetic acid, is a liquid with a density of 1.049 g/mL. Which calculation correctly shows how to determine the volume of glacial acetic acid necessary to prepare 250 mL of 0.400 M CH3CO2H(aq)?
A.
B.
C.
D.
E.
17. Pure acetic acid, CH3CO2H(l), has a density of 1.049 g/mL. To prepare 1.00 L of 6.00 M CH3CO2H(aq), one may A. dilute 175 g of acetic acid to a volume of 1.00 L. B. dilute 343 mL of acetic acid to a volume of 1.00 L. C. dilute 360 mL of acetic acid to a volume of 1.00 L. D. dilute 382 mL of acetic acid to a volume of 1.00 L. E. dilute 1049 g of acetic acid to a volume of 1.00 L.
18. A 15 meter by 12 meter pool of water has a depth of 2.2 meters. What mass of silver ion is present in the reservoir if the concentration of silver ion is 0.14 ppm? (1 m3 = 1000 L; assume the density of the solution is 1.00 g/mL) A. 5.5 ´ 10–4 g B. 5.5 ´ 10–2 g C. 0.55 g D. 5.5 g E. 55 g
19. If sea water contains 15 ppm gold, how many kilograms of sea water must be processed to remove 1.00 g of gold? A. 67 kg B. 97 kg C. 150 kg D. 6.7 ´ 103 kg E. 15 ´ 104 kg
20. Silver chloride is a relatively insoluble salt. Only 1.92 mg of AgCl will dissolve per liter of water at 25C. How many parts per million of Ag+ can be present in water at 25C? Assume the density of the solution equals the density of water, 1.00 g/mL. A. 0.52 ppm B. 0.93 ppm C. 1.45 ppm D. 9.30 ppm E. 52.0 ppm
21. Silver chloride is a relatively insoluble salt. Only 1.92 mg of AgCl will dissolve per liter of water at 25C. What concentration of Ag+, in molarity units, can be present in water at 25C? A. 3.35 ´ 10-6 M B. 6.70 ´ 10-6 M C. 1.34 ´ 10-5 M D. 9.60 ´ 10-4 M E. 1.92 ´ 10-3 M
22. What concentration of silver nitrate (in ppm) is present in 7.1 ´ 10–7 M AgNO3(aq)? For very dilute aqueous solutions, you can assume the solution's density is 1.0 g/mL. The molar mass of AgNO3 is 169.9 g/mol. A. 0.0071 ppm B. 0.12 ppm C. 0.71 ppm D. 1.7 ppm E. 8.3 ppm
23. A gas mixture has mole fractions of 0.24 oxygen and 0.76 nitrogen. If the total pressure of the gases is 1.44 atm at 325 K, what is the concentration, in molarity, of oxygen? (R = 0.0821 Latm/molK) A. 0.013 M B. 0.017M C. 0.041 M D. 0.54 M E. 0.069 M
24. If 77.5 g of ethylene glycol (HOCH2CH2OH) is added to 422.5 g of water, what is the mole fraction of ethylene glycol? A. 0.00296 B. 0.0506 C. 0.183 D. 0.949 E. 2.96
25. Arrange the molecules below in order of increasing solubility in water.
A. dimethyl ether < methane < methanol B. dimethyl ether < methanol < methane C. methane < methanol< dimethyl ether D. methanol < dimethyl ether < methane E. methane < dimethyl ether < methanol
26. Which of the following liquids will be miscible with water in any proportions: ethanol (CH3CH2OH), carbon tetrachloride (CCl4), hexane (C6H14), and/or formic acid (HCO2H)? A. ethanol and carbon tetrachloride B. carbon tetrachloride and hexane C. ethanol and formic acid D. ethanol, carbon tetrachloride, and benzene E. carbon tetrachloride, and formic acid
27. Arrange the molecules below in order of increasing solubility in water.
A. methane < methanol < chloromethane B. methane < chloromethane < methanol C. chloromethane < methanol < methane D. chloromethane < methane < methanol E. methanol < methane < chloromethane
28. All of the following statements are correct EXCEPT A. the solubility of a gas in water decreases as the water temperature increases. B. dissolving a solid in water is usually an exothermic process. C. when an equilibrium is established between molecules in a solid and a solution, the solution is said to be saturated. D. if a precipitate forms when a solution is cooled, the solution is supersaturated. E. network covalent solids are usually insoluble in water.
29. A substance that dissolves in water and conducts electricity when present in an aqueous solution is A. metallic. B. ionic. C. molecular. D. network covalent. E. none of the above.
30. Which of the following statements concerning the solubility of a gas in a liquid are true? 1. 2. 3.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1 and 3
Solubility decreases with increasing temperature. Solubility increases as the pressure of the gas over the liquid increases. Solubility is dependent on the surface area of the the liquid.
31. Henry's law states that gas solubility is A. directly proportional to temperature of the solution. B. directly proportional to the molar mass of the gas. C. inversely proportional to the combined pressure of all gases over the solution. D. inversely proportional to the pressure of the gas over the solution. E. directly proportional to the pressure of the gas over the solution.
32. The Henry's law constant for O2 in water at 25 C is 1.26 ´ 10–3 M/atm. What partial pressure of O2 is necessary to achieve an equilibrium concentration of 1.5 ´ 10–3 M O2? A. 1.9 ´ 10–6 atm B. 0.24 atm C. 0.84 atm D. 1.2 atm E. 5.3 ´ 105 atm
33. The Henry's law constant for the solubility of nitrogen in water is 6.4 ´ 10-4 M/atm at 25C. At 0.75 atm of N2, what mass of N2(g) dissolves in 1.0 L of water at 25C? A. 4.8 ´ 10-4 g B. 8.5 ´ 10-4 g C. 4.5 ´ 10-3 g D. 1.3 ´ 10-2 g E. 2.4 ´ 10-2 g
34. In 1.00 atm of pure oxygen, the solubility of O2(g) in water is 1.26 ´ 10-3 M at 25.0C. The mole fraction of oxygen in air is 0.210. If the atmospheric pressure is 0.979 atm, what is the solubility of oxygen in air at 25.0C? A. 2.59 ´ 10-4 M B. 2.65 ´ 10-3 M C. 5.87 ´ 10-3 M D. 6.00 ´ 10-3 M E. 3.86 ´ 103 M
35. At 1.00 atm, 1.64 ´ 10-3 g H2(g) will dissolve in 1.0 L of water. What pressure of gas is necessary to obtain a concentration of 1.0 ´ 10-3 M H2(g)? A. 6.1 ´ 10-4 atm B. 4.4 ´ 10-2 atm C. 1.2 atm D. 2.2 ´ 103 atm E. 1.6 ´ 103 atm
36. All of the following are colligative properties EXCEPT A. gas solubility as a function of partial pressure above a solution (Henry's law). B. osmotic pressure. C. vapor pressure lowering. D. boiling point elevation. E. freezing point depression.
37. An aqueous solution contains 11.5 g of NaCl in 250.2 g of water. Calculate the vapor pressure of this solution at 25.0C. The vapor pressure of pure water is 23.8 mm Hg at 25.0C. A. 3.38 mm Hg B. 21.8 mm Hg C. 23.5 mm Hg D. 24.1 mm Hg E. 31.3 mm Hg
38. What is the equilibrium partial pressure of water vapor above a mixture of 44.0 g H2O and 56.0 g HOCH2CH2OH at 35C. The partial pressure of pure water at 35.0C is 42.2 mm Hg. Assume ideal behavior for the solution. A. 0.730 Hg B. 18.7 mm Hg C. 23.6 mm Hg D. 30.8 mm Hg E. 58.8 mm Hg
39. What mass of ethylene glycol, when mixed with 225 g H2O, will reduce the equilibrium vapor pressure of H2O from 1.00 atm to 0.800 atm at 100C? The molar masses of water and ethylene glycol are 18.02 g/mol and 62.07 g/mol, respectively. Assume ideal behavior for the solution. A. 15.6 g B. 49.9 g C. 194 g D. 969 g E. 3.10 ´ 103 g
40. Which of the following aqueous solutions will freeze at the lowest temperature? A. 0.10 m KCl B. 0.20 m C6H12O6 (glucose) C. 0.050 m AlCl3 D. 0.15 m SrBr2 E. All of the above freeze at the same temperature.
41. What concentration unit is necessary for the calculation of boiling point elevation? A. molarity B. molality C. mass fraction D. mole fraction E. parts per million
42. What concentration unit is necessary for the calculation of vapor pressure lowering? A. molarity B. molality C. mass fraction D. mole fraction E. density
43. The molal boiling point constant for water is 0.52C/m. At what temperature will a mixture of 45.0 g of NaCl and 0.500 kg of water boil? A. 98.4C B. 99.2C C. 100.0C D. 100.8C E. 101.6C
44. What is the boiling point of a solution containing 0.80 g caffeine, C8H10N4O2, dissolved in 13.20 g benzene? The boiling point of pure benzene is 80.1C and the molal boiling point constant, Kb, is 2.53C/m. A. 79.8C B. 80.4C C. 80.9C D. 85.2C E. 88.2C
45. The molal freezing point constant for water is –1.86C/m. At what temperature will a solution containing 8.27 g CaCl2 and 45.0 g H2O begin to freeze? Assume that no ion-pairing occurs between Ca2+ and Cl–. A. –9.24C B. –4.62C C. –0.804C D. –0.749C E. +4.62C
46. What is the molar mass of a nonpolar molecular compound if 3.42 grams dissolved in 41.8 grams benzene begins to freeze at 1.17C? The freezing point of pure benzene is 5.50C and the molal freezing point constant, Kf, is –5.12C/m. A. 2.89 g/mol B. 69.2 g/mol C. 96.7 g/mol D. 126 g/mol E. 358 g/mol
47. Equal masses of water and ethylene glycol (C2H6O2) are mixed. At what temperature will the mixture freeze? The molal freezing point constant for water is -1.86C/m. A. -115C B. -93.0C C. -42.0C D. -30.0C E. -0.93C
48. Which of the following statements concerning osmosis is/are CORRECT? 1. 2. 3.
Osmosis involves the movement of a solvent through a semipermeable membrane. Solvents move from regions of high solute concentration to regions of lower solute concentration. Osmotic pressure is a colligative property.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 49. At 25C, what is the osmotic pressure of a homogeneous solution consisting of 18.0 g urea (CON2H4) diluted with water to 3.00 L? (R = 0.0821 Latm/molK) A. 0.205 atm B. 2.44 atm C. 7.33 atm D. 12.3 atm E. 14.7 atm
50. An aqueous solution is composed of 7.50 g NaCl (MM = 58.44 g/mol) diluted to 0.100 L. Calculate the osmotic pressure of the solution at 298 K. (R = 0.0821 Latm/molK) A. 5.83 atm B. 9.22 atm C. 18.3 atm D. 31.4 atm E. 62.8 atm
51. A solution is prepared by dissolving 5.00 g of an unknown molecular solid in water to make 1.00 L of solution. The osmotic pressure of the solution is 1.61 atm at 25C. What is the molar mass of the solute? (R = 0.0821 Latm/molK) A. 6.37 g/mol B. 58.44 g/mol C. 76.0 g/mol D. 102 g/mol E. 180.2 g/mol
52. A solution is prepared by dissolving 4.78 g of an unknown nonelectrolyte in enough water to make 375 mL of solution. The osmotic pressure of the solution is 1.33 atm at 27C. What is the molar mass of the solute? (R = 0.0821 Latm/molK) A. 0.0203 g/mol B. 21.2 g/mol C. 49.4 g/mol D. 96.8 g/mol E. 236 g/mol
53. A solution is prepared by dissolving 4.21 g of a nonelectrolyte in 50.0 g of water. If the boiling point increases by 0.203C, what is the molar mass of the solute? The boiling point elevation constant for water is 0.512C/m. A. 33.4 g/mol B. 111 g/mol C. 172 g/mol D. 212 g/mol E. 810 g/mol
54. Which of the following electrolytes is likely to have a van't Hoff factor equal to 3? A. CaI2 B. Na3PO4 C. KCl D. answers a and b E. answers a, b, and c
55. A 0.230 m solution of an unknown electrolyte depresses the freezing point of water by 0.821C. What is the Van't Hoff factor for this electrolyte? The freezing point depression constant for water is 1.86C/m. A. 0.521 B. 1.92 C. 2.00 D. 2.30 E. 4.41
56. Maple syrup is made from the sap of the maple tree. When the sap is tapped from the maple tree it is 2.0% by mass sucrose. Maple syrup is 66% by mass sucrose. What mass of sap would be required to make 0.5 kg of maple syrup? A. 0.50 kg B. 2.0 kg C. 66 kg D. 15 kg E. 17 kg
Chapter 11--Rate of Reaction 1. Which statement concerning relative rates of reaction is correct for the chemical equation given below? 2CH3OH(g) + 3O2(g) ® 2CO2(g) + 4H2O(g) A. The rate of disappearance of CH3OH is equal to the rate of disapperance of O2. B. The rate of disappearance of CH3OH is two times the rate of appearance of H2O. C. The rate of disappearance of CH3OH is half the rate of appearance of CO2. D. The rate of appearance of H2O is two times the rate of appearance of CO2. E. The rate of appearance of H2O is four times the rate of disappearance of CH3OH.
2. Express the rate of the reaction below in terms of the formation of dinitrogen pentaoxide. 4NO2(g) + O2(g) ® 2N2O5(g)
A. B.
C. D.
E.
3. For the reaction below relate the rate of disappearance of hydrogen to the rate of formation of ammonia. N2(g) + 3H2(g) ® 2NH3(g)
A.
B.
C.
D.
E.
4. The rate of reaction for the formation of carbon monoxide is measured at 1.24 mol/Lhr. What is the rate of formation of carbon monoxide in units of mol/Ls? CH3CHO(g) ® CH4(g) + CO(g) A. 3.44 ´ 10-4 mol/Ls B. 2.07 ´ 10-2 mol/Ls C. 1.24 mol/Ls D. 74.4 mol/Ls E. 4.64 ´ 103 mol/Ls
5. For the reaction below, if the rate of appearance of Br2 is 0.180 mol/Ls, what is the rate of disappearance of NOBr? 2NOBr(g) ® 2NO(g) + Br2(g) A. –0.360 mol/Ls B. –0.090 mol/Ls C. 0.090 mol/Ls D. 0.180 mol/Ls E. 0.360 mol/Ls
6. Assume the reaction below 2NO(g) + O2(g) ® 2NO2(g) proceeds via the following rate expression:
Which of the following statements concerning the above chemical reaction and rate equation is/are CORRECT? 1. 2. 3.
The reaction is second-order with respect to NO. The rate of disappearance of O2 is two times the rate of appearance of NO2. According to the balanced chemical equation, the reaction is fifth-order overall.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 2 and 3 7. Dinitrogen pentaoxide decomposes to nitrogen dioxide and oxygen according to the following balanced chemical equation and rate expression. 2N2O5(g) ® 4NO2(g) + O2(g) rate = k[N2O5] What is the overall reaction order? A. 0 B. 1 C. 2 D. 5 E. 7
8. Determine the overall reaction order for the chemical expression and rate expression below. 2NO(g) + Cl2(g) ® 2NOCl(g) rate = k[NO]2 ´ [Cl2] A. 1 B. 2 C. 3 D. 4 E. 5
9. What are the units of the rate constant for the rate expression for the following chemical equation? 2NO(g) + Cl2(g) ® 2NOCl(g) rate = k[NO]2 ´ [Cl2] A. mol/Ls B. mol2/L2s C. L/mols D. L2/mol2s E. L3/mol3s
10. Given the initial rate data for the decomposition reaction, A®B+C determine the rate expression for the reaction.
[A] (mol/L) 0.084 0.063 0.042
–D[A]/Dt (mol/Ls) 12.4 ´ 10–6 9.3 ´ 10–6 6.2 ´ 10–6
A.
B.
C.
D.
E. 11. The reaction rate of CO and NO2 in the reaction CO(g) + NO2(g) ® CO2(g) + NO(g) is measured using the initial rates method. The results are tabulated below.
[CO] (mol/L) 8.00 ´ 10-4 8.00 ´ 10-4 1.60 ´ 10-3
NO2 (mol/L) 5.50 ´ 10 -4 1.10 ´ 10-3 5.50 ´ 10-4
-D[CO]/Dt (mol/Ls) 8.40 ´ 10-8 1.68 ´ 10-7 1.68 ´ 10-7
Determine the rate expression and calculate the rate constant for the reaction.
A. rate = 0.191[CO] ´ [NO2] B. rate = 0.191[CO]2 ´ [NO2] C. rate = 0.191[CO] ´ [NO2]2 D. rate = 5.24[CO] ´ [NO2] E. rate = 5.24[CO]2 ´ [NO2] 12. Given the initial rate data for the reaction A + B ® C, determine the rate expression for the reaction.
[A], (mol/L) 0.0344 0.0516 0.0344
[B], (mol/L) 0.160 0.160 0.272
D[C]/Dt (mol/Ls) 2.11 ´ 10–2 3.17 ´ 10–2 6.10 ´ 10–2
A.
B.
C.
D.
E. 13. The initial rates method was used to study the reaction below. A + 3B ® 2C
[A] (mol/L) 0.210 0.210 0.420
[B] (mol/L) 0.150 0.300 0.300
Determine the rate expression and calculate the rate constant for the reaction.
A. rate = 0.515[A] ´ [B] B. rate = 0.515[A]2 ´ [B] C. rate = 0.721[A]2 ´ [B] D. rate = 0.721[A] ´ [B]2 E. rate = 0.721[A]2 ´ [B]2
-D[A]/Dt (mol/Ls) 3.41 ´ 10-3 1.36 ´ 10-2 2.73 ´ 10 -2
14. The initial rates method was used to study the reaction below. 2A + B ® C
[A] (mol/L) 0.555 0.555 0.278
[B] (mol/L) 0.300 0.150 0.150
D[C]/Dt (mol/Ls) 7.11 ´ 10-4 3.56 ´ 10-4 3.56 ´ 10 -4
Determine the rate expression and calculate the rate constant for the reaction.
A. rate = 7.90 ´ 10-3 [A] ´ [B]2 B. rate = 2.37 ´ 10-3 [A] ´ [B]2 C. rate = 2.37 ´ 10-3 [B] D. rate = 2.37 ´ 10-3 [B]2 E. rate = 7.90 ´ 10-3 [B]2 15. The initial rates method was used to study the reaction below. 2A + B + C ® D + E
[A] (mol/L) 0.150 0.150 0.150 0.300
[B] (mol/L) 0.250 0.125 0.250 0.125
Determine the rate expression and calculate the rate constant for the reaction.
A. rate = 8.71 ´ 10-4[A]2 ´ [B] ´ [C] B. rate = 5.23 ´ 10-4[A] ´ [B]2 ´ [C] C. rate = 2.90 ´ 10-3[A]2 ´ [B] ´ [C]2 D. rate = 8.71 ´ 10-4[A] ´ [B] ´ [C]2 E. rate = 1.31 ´ 10-4[A] ´ [B] ´ [C]
[C] (mol/L) 0.300 0.300 0.600 0.300
D[D]/Dt (mol/Ls) 1.47 ´ 10-6 3.68 ´ 10-7 2.94 ´ 10-6 7.35 ´ 10-7
16. For the reaction A + 2B ® C, the rate law is
. What are the units of the rate constant where time is measured in seconds? A. B. C.
D. E.
17. The reaction of NO and O2 produces NO2. 2 NO(g) + O2(g) ® 2 NO2(g) The reaction is second-order with respect to NO(g) and first-order with respect to O2(g). At a given temperature, the rate constant, k, equals 4.7 ´ 102 M–2s–1. What is the rate of reaction when the initial concentrations of NO and O2 are 0.025 M and 0.015 M, respectively? A. 2.6 ´ 10–3 M/s B. 4.4 ´ 10–3 M/s C. 0.18 M/s D. 2.0 ´ 10–8 M/s E. 3.8 ´ 102 M/s
18. For the second-order decomposition reaction, 2A ® B which of the following relationships yields a straight line plot?
A. ln [A] versus k B. ln [A]2 versus t C. ln [A] versus t D. 1/[A] versus t E. 1/[A] versus k
rate = k[A]2
19. Which of the following expressions corresponds to the integrated rate equation for a first-order decomposition reaction? A. [A] = -kt B. [A] = [A]0 -kt C. ln [A] = ln [A]0 -kt D. E.
20. The rate constant of a first-order decomposition reaction is 0.0147 s–1. If the initial concentration of reactant is 0.178 M, what is the concentration of reactant after 30.0 seconds? A. 8.72 ´ 105 M B. 0.0645 M C. 0.115 M D. 0.0785 M E. 0.643 M
21. The rate constant for the decomposition of cyclobutane is 2.08 ´ 10-2 s-1 at high temperatures. C4H8(g) ® 2C2H4(g) How many seconds are required for an initial concentration of 0.100 M C4H8(g) to decrease to 0.0450 M? A. 0.00114 s B. 1.07 s C. 2.64 s D. 38.4 s E. 874 s
22. At a high temperature, the first-order decomposition of N2O5(g) produces NO2(g) and O2(g). If the initial concentration of 0.400 M N2O5(g) is reduced to 0.169 M after 118 seconds, what is the rate constant for the reaction? A. 1.96 ´ 10-3 s-1 B. 2.29 ´ 10-3 s-1 C. 4.37 ´ 10-3 s-1 D. 7.30 ´ 10-3 s-1 E. 1.37 ´ 102 s-1
23. The reaction A ® B follows first-order kinetics with a half-life of 21.7 hours. If the concentration of A is 0.023 M after 48.0 hours, what is the initial concentration of A? A. 0.0050 M B. 0.051 M C. 0.51 M D. 0.11 M E. 2.0 ´ 102 M
24. For the first-order reaction below, the concentration of product B after 24.2 seconds is 0.322 M. If k = 8.75 ´ 10-2 s-1, what was the initial concentration of A? A ® 2B
rate = k[A]
A. 0.0341 M B. 0.183 M C. 1.34 M D. 2.68 M E. 29.3 M 25. Which equation is used to calculate the half-life of the first-order equation below? A ® 2B
rate = k[A]
A. t1/2 = 2 ´ ln k B. t1/2 = (ln k)/[A]o C. t1/2 = ln (k/2) D. t1/2 = ln ([A]o/k) E. t1/2 = (ln 2)/k 26. Which equation is used to calculate the half-life of the second-order equation below? 2A ® B
rate = k[A]2
A.
B.
C.
D. E. 27. Which equation is used to calculate the half-life of the zero-order equation below? A ® 2B
rate = k
A.
B.
C. D. E. t1/2 = k 28. For the second-order reaction below, the rate constant of the reaction is 9.4 ´ 10–3 M–1s–1. How long (in seconds) is required to decrease the concentration of A from 2.16 M to 0.40 M? 2A ® B
A. 2.0 ´ 101 s B. 7.8 ´ 101 s C. 1.8 ´ 102 s D. 1.9 ´ 102 s E. 2.2 ´ 102 s
rate = k[A]2
29. For the second-order reaction below, the initial concentration of A is 1.34 M. If k = 7.61 ´ 10-4 s-1, what is the concentration of A after 18.3 minutes? 2A ® B
rate = k[A]2
A. 0.504 M B. 0.581 M C. 0.632 M D. 0.836 M E. 1.12 M 30. For the second-order reaction below, the initial concentration of A is 1.00 M. If the concentration of A is reduced to 0.43 M after 75 seconds, what is the rate constant? 2A ® B
rate = k[A]2
A. 5.2 ´ 10-4 L/mols B. 5.7 ´ 10-3 L/mols C. 7.6 ´ 10-3 L/mols D. 1.1 ´ 10-2 L/mols E. 1.8 ´ 10-2 L/mols 31. For the second-order reaction below, the concentration of product B after 132 seconds is 0.0281 M. If the initial concentration of A is 0.932 M, what is the rate constant? 2A ® B
rate = k[A]2
A. 1.43 ´ 10-5 L/mols B. 4.71 ´ 10-4 L/mols C. 5.22 ´ 10-4 L/mols D. 6.63 ´ 10-3 L/mols E. 96.4 L/mols 32. What is the half-life of a first-order reaction if the rate constant is 6.2 ´ 10–3 s–1? A. 8.9 ´ 10–3 s B. 0.097 s C. 5.77 s D. 1.0 ´ 102 s E. 1.6 ´ 102 s
33. For the first-order reaction below, the initial concentration of A is 0.240 M. If the concentration of A decreases to 0.0800 M after 21.8 hours, what is the half-life of the reaction? A®B
rate = k[A]
A. 0.0504 hrs B. 1.28 hrs C. 7.26 hrs D. 10.4 hrs E. 13.8 hrs 34. Hydrogen peroxide decomposes into water and oxygen in a first-order process. H2O2(aq) ® H2O(l) + 1/2 O2(g) At 20.0C, the half-life for the reaction is 3.92 ´ 104 seconds. If the initial concentration of hydrogen peroxide is 0.52 M, what is the concentration after 7.00 days? A. 1.2 ´ 10-5 M B. 0.034 M C. 0.074 M D. 0.22 M E. 0.52 M
35. A first-order reaction has a half-life of 4.54 seconds. How much time is required for the reactant to be reduced to 6.25% of its initial concentration? A. 1.14 s B. 9.08 s C. 13.6 s D. 18.2 s E. 93.6 s
36. For the first-order reaction below, the initial concentration of A is 0.80 M. What is the half-life of the reaction if the concentration of A decreases to 0.10 M in 54 seconds? A®B
A. 18 s B. 24 s C. 36 s D. 48 s E. 51 s
rate = k[A]
37. A first-order reaction has a half-life of 2.10 hours. How much time is required for the reactant to be reduced to 33.0% of its initial concentration? A. 2.45 hours B. 2.89 hours C. 3.36 hours D. 3.93 hours E. 4.21 hours
38. The decomposition of formic acid follows first-order kinetics. HCO2H(g) ® CO2(g) + H2(g) The half-life for the reaction at 550C is 24 seconds. How many seconds does it take for the formic acid concentration to decrease by 87.5%? A. 24 s B. 36 s C. 48 s D. 72 s E. 96 s
39. Which of the following factors are likely to affect the rate of a chemical reaction? 1. 2. 3.
the presence of a catalyst the temperature of the reactants the physical state (solid, liquid, or gas) of the reactants
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 40. The effect of a catalyst on a chemical reaction is to A. enable an alternate path for the reaction that has a lower activation barrier. B. increase the energy of the products. C. increase the activation barrier for the forward reaction. D. increase the energy of the reactants. E. increase the frequency of collisions between reactants and products.
41. The transition-state model for reaction rate assumes that the activated complex 1. 2. 3.
is in equilibrium with the reactants. is gaseous at low concentrations. may decompose to form products or may revert to reactants.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 42. The correct form of the Arrhenius equation is A. Ea = Ae-k/RT B. C. D. Ea = Aek/RT E. A = Eae-k/RT
43. For a given reaction, the rate constant doubles when the temperature is increased from 45.0C to 73.0C. What is the activation energy for this reaction? (R = 8.31 J/Kmol) A. 0.676 kJ/mol B. 9.85 kJ/mol C. 16.1 kJ/mol D. 22.7 kJ/mol E. 65.4 kJ/mol
44. For a given reaction, the activation energy is 63.9 kJ/mol. If the reaction rate constant is 4.1 ´ 10–3 L/mols at 32.0C, what is the reaction rate constant at –5.0C? (R = 8.31 J/Kmol) A. 1.2 ´ 10–3 L/mols B. 1.4 ´ 10–6 L/mols C. 1.3 ´ 10–4 L/mols D. 4.0 ´ 10–3 L/mols E. 1.3 ´ 10–1 L/mols
45. The rate constant at 373 K for a certain reaction is 8.29 ´ 10-4 s-1 and the activation energy is 12.0 kJ/mole. What is the value of the constant, A, in the Arrhenius equation? A. 0.00832 s-1 B. 0.0398 s-1 C. 0.120 s-1 D. 0.998 s-1 E. 25.1 s-1
46. For a certain reaction, the activation energy is 52.1 kJ/mole. By what ratio will the rate constant change if the temperature is decreased from 175C to 75C? A. 0.00402 B. 0.0179 C. 0.996 D. 121 E. 249
47. All of the following statements are correct EXCEPT A. a heterogeneous catalyst is in a different phase from the reaction mixture. B. enzymes are protein molecules that catalyze reactions. C. catalysts are not consumed in reactions. D. a reaction mechanism describes the path of a reaction at the molecular level. E. elementary steps in a reaction mechanism are always unimolecular.
48. The elementary steps for the catalyzed decomposition of dinitrogen monoxide are shown below. N2O(g) + NO(g) ® N2(g) + NO2(g) NO2(g) ® NO(g) + 1/2O2(g) Which of the following statements is/are CORRECT? 1. 2. 3.
The overall balanced reaction is N2O(g) ® N2(g) + 1/2O2(g). NO2(g) is a catalyst for the reaction. NO(g) is a reaction intermediate.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 49. The elementary steps for the catalyzed decomposition for dinitrogen monoxide are shown below. Identify the catalyst. NO(g) + N2O(g) ® N2(g) + NO2(g) 2NO2(g) ® 2NO(g) + O2(g) A. NO(g) B. N2O(g) C. N2(g) D. NO2(g) E. O2(g)
50. The elementary steps for a catalyzed reaction are shown below. Identify the catalyst. Identify the reactive intermediate. H2O2(aq) + I-(aq) ® H2O(l) + IO-(aq) IO-(aq) + H2O2(aq) ® H2O(l) + O2(g) + I-(aq) A. The catalyst is H2O(l); the reactive intermediate is I-(aq). B. The catalyst is IO-(aq); the reactive intermediate is I-(aq). C. The catalyst is I-(aq); the reactive intermediate is H2O2(aq). D. The catalyst is I-(aq); the reactive intermediate is IO-(aq). E. The catalyst is H2O2(aq); the reactive intermediate is I-(aq).
51. A possible mechanism for the decomposition of ozone to oxygen in the atmosphere is O3(g) O2(g) + O O + O3(g) ® 2O2(g)
(fast equilibrium) (slow)
What is a rate law that is consistent with this mechanism?
A. rate = k[O3] B. rate = k[O3]2 C. rate = k[O3] ´ [O] D. rate = k[O3]2 ´ [O2] E. rate = k[O3]2 ´ [O2]-1 52. A possible reaction mechanism for the reaction of nitrogen dioxide with carbon monoxide is 2NO2(g) ® NO3(g) + NO(g) NO 3(g) + CO(g) ® NO2(g) + CO2(g)
(slow) (fast)
What is the overall reaction and the most probable rate law for the reaction?
A. NO2(g) + CO(g) ® NO(g) + CO2(g); rate = k[NO2] B. NO2(g) + CO(g) ® NO(g) + CO2(g); rate = k[NO2]2 C. NO2(g) + CO(g) ® NO(g) + CO2(g); rate = k[NO3]´[CO] D. 2NO2(g) + CO(g) ® NO(g) + CO2(g); rate = k[NO2]´[ NO3]´[CO] E. 2NO2(g) + CO(g) ® NO(g) + CO2(g); rate = k[NO3]´[CO] 53. Nitrogen dioxide reacts with carbon monoxide to produce nitrogen monoxide and carbon dioxide. NO2(g) + CO(g) ® NO(g) + CO2(g) A proposed mechanism for this reaction is 2NO2(g) NO3(g) + NO(g) NO3(g) + CO(g) ® NO2(g) + CO2(g) What is a rate law that is consistent with the proposed mechanism?
A. rate = k[NO2]2´[CO]´ [NO]–1 B. rate = k[NO2]2´[CO] C. rate = k[NO2]´[CO] D. rate = k[NO3]´[CO] E. rate = k[NO2]2
(fast, equilibrium) (slow)
54. For the overall reaction A + 2B ® C which of the following mechanisms yields the correct overall chemical equation and is consistent with the rate equation below? rate = k[A]´[B] A. A + B I (fast) I+A®C (slow) B. A + B ® I (slow) I+B®C (fast) C. 2B ® I (slow) A+I®C (fast) D. 2B I (fast) I+A®C (slow) E. A + 2B I (fast) I+B®C+B (slow)
55. For the overall reaction 2A + B ® C which of the following mechanisms are consistent with a rate equation of rate = k[A]2´[B]? A. A + B I (fast) I + A ® C (slow) B. A + B ® I (slow) I + A ® C (fast) C. 2A ® I (slow) B + I ® C (fast) D. 2A I (fast) I + B ® C (slow) E. Answers a and d are both correct.
56. What did Henry Eyring develop that changed how scientists study rates of reactions? A. the catalyst model B. the transition-state model C. the liquid-state model D. the temperature dependence model E. the enzyme model
57. Ozone in the upper atmosphere is important to life on earth because it converts harmful UV radiation to heat. How does the release of CFCs into the atmosphere affect the concentration of ozone in the upper atmosphere? A. the Cl in the CFC catalyzes the decomposition of ozone decreasing its concentration B. the Cl in the CFC catalyzes the production of ozone increasing its concentration C. the NO2 in the CFC catalyzes the decomposition of ozone decreasing its concentration D. the NO2 in the CFC catalyzes the production of ozone increasing its concentration E. the CFC's change the ozone into photochemical smog
Chapter 11--Rate of Reaction Key
1. Which statement concerning relative rates of reaction is correct for the chemical equation given below? 2CH3OH(g) + 3O2(g) ® 2CO2(g) + 4H2O(g) A. The rate of disappearance of CH3OH is equal to the rate of disapperance of O2. B. The rate of disappearance of CH3OH is two times the rate of appearance of H2O. C. The rate of disappearance of CH3OH is half the rate of appearance of CO2. D. The rate of appearance of H2O is two times the rate of appearance of CO2. E. The rate of appearance of H2O is four times the rate of disappearance of CH3OH.
2. Express the rate of the reaction below in terms of the formation of dinitrogen pentaoxide. 4NO2(g) + O2(g) ® 2N2O5(g)
A. B.
C. D.
E.
3. For the reaction below relate the rate of disappearance of hydrogen to the rate of formation of ammonia. N2(g) + 3H2(g) ® 2NH3(g)
A.
B.
C.
D.
E.
4. The rate of reaction for the formation of carbon monoxide is measured at 1.24 mol/Lhr. What is the rate of formation of carbon monoxide in units of mol/Ls? CH3CHO(g) ® CH4(g) + CO(g) A. 3.44 ´ 10-4 mol/Ls B. 2.07 ´ 10-2 mol/Ls C. 1.24 mol/Ls D. 74.4 mol/Ls E. 4.64 ´ 103 mol/Ls
5. For the reaction below, if the rate of appearance of Br2 is 0.180 mol/Ls, what is the rate of disappearance of NOBr? 2NOBr(g) ® 2NO(g) + Br2(g) A. –0.360 mol/Ls B. –0.090 mol/Ls C. 0.090 mol/Ls D. 0.180 mol/Ls E. 0.360 mol/Ls
6. Assume the reaction below 2NO(g) + O2(g) ® 2NO2(g) proceeds via the following rate expression:
Which of the following statements concerning the above chemical reaction and rate equation is/are CORRECT? 1. 2. 3.
The reaction is second-order with respect to NO. The rate of disappearance of O2 is two times the rate of appearance of NO2. According to the balanced chemical equation, the reaction is fifth-order overall.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 2 and 3 7. Dinitrogen pentaoxide decomposes to nitrogen dioxide and oxygen according to the following balanced chemical equation and rate expression. 2N2O5(g) ® 4NO2(g) + O2(g) rate = k[N2O5] What is the overall reaction order? A. 0 B. 1 C. 2 D. 5 E. 7
8. Determine the overall reaction order for the chemical expression and rate expression below. 2NO(g) + Cl2(g) ® 2NOCl(g) rate = k[NO]2 ´ [Cl2] A. 1 B. 2 C. 3 D. 4 E. 5
9. What are the units of the rate constant for the rate expression for the following chemical equation? 2NO(g) + Cl2(g) ® 2NOCl(g) rate = k[NO]2 ´ [Cl2] A. mol/Ls B. mol2/L2s C. L/mols D. L2/mol2s E. L3/mol3s
10. Given the initial rate data for the decomposition reaction, A®B+C determine the rate expression for the reaction.
[A] (mol/L) 0.084 0.063 0.042
–D[A]/Dt (mol/Ls) 12.4 ´ 10–6 9.3 ´ 10–6 6.2 ´ 10–6
A.
B.
C.
D.
E. 11. The reaction rate of CO and NO2 in the reaction CO(g) + NO2(g) ® CO2(g) + NO(g) is measured using the initial rates method. The results are tabulated below.
[CO] (mol/L) 8.00 ´ 10-4 8.00 ´ 10-4 1.60 ´ 10-3
NO2 (mol/L) 5.50 ´ 10 -4 1.10 ´ 10-3 5.50 ´ 10-4
-D[CO]/Dt (mol/Ls) 8.40 ´ 10-8 1.68 ´ 10-7 1.68 ´ 10-7
Determine the rate expression and calculate the rate constant for the reaction.
A. rate = 0.191[CO] ´ [NO2] B. rate = 0.191[CO]2 ´ [NO2] C. rate = 0.191[CO] ´ [NO2]2 D. rate = 5.24[CO] ´ [NO2] E. rate = 5.24[CO]2 ´ [NO2] 12. Given the initial rate data for the reaction A + B ® C, determine the rate expression for the reaction.
[A], (mol/L) 0.0344 0.0516 0.0344
[B], (mol/L) 0.160 0.160 0.272
D[C]/Dt (mol/Ls) 2.11 ´ 10–2 3.17 ´ 10–2 6.10 ´ 10–2
A.
B.
C.
D.
E. 13. The initial rates method was used to study the reaction below. A + 3B ® 2C
[A] (mol/L) 0.210 0.210 0.420
[B] (mol/L) 0.150 0.300 0.300
Determine the rate expression and calculate the rate constant for the reaction.
A. rate = 0.515[A] ´ [B] B. rate = 0.515[A]2 ´ [B] C. rate = 0.721[A]2 ´ [B] D. rate = 0.721[A] ´ [B]2 E. rate = 0.721[A]2 ´ [B]2
-D[A]/Dt (mol/Ls) 3.41 ´ 10-3 1.36 ´ 10-2 2.73 ´ 10 -2
14. The initial rates method was used to study the reaction below. 2A + B ® C
[A] (mol/L) 0.555 0.555 0.278
[B] (mol/L) 0.300 0.150 0.150
D[C]/Dt (mol/Ls) 7.11 ´ 10-4 3.56 ´ 10-4 3.56 ´ 10 -4
Determine the rate expression and calculate the rate constant for the reaction.
A. rate = 7.90 ´ 10-3 [A] ´ [B]2 B. rate = 2.37 ´ 10-3 [A] ´ [B]2 C. rate = 2.37 ´ 10-3 [B] D. rate = 2.37 ´ 10-3 [B]2 E. rate = 7.90 ´ 10-3 [B]2 15. The initial rates method was used to study the reaction below. 2A + B + C ® D + E
[A] (mol/L) 0.150 0.150 0.150 0.300
[B] (mol/L) 0.250 0.125 0.250 0.125
Determine the rate expression and calculate the rate constant for the reaction.
A. rate = 8.71 ´ 10-4[A]2 ´ [B] ´ [C] B. rate = 5.23 ´ 10-4[A] ´ [B]2 ´ [C] C. rate = 2.90 ´ 10-3[A]2 ´ [B] ´ [C]2 D. rate = 8.71 ´ 10-4[A] ´ [B] ´ [C]2 E. rate = 1.31 ´ 10-4[A] ´ [B] ´ [C]
[C] (mol/L) 0.300 0.300 0.600 0.300
D[D]/Dt (mol/Ls) 1.47 ´ 10-6 3.68 ´ 10-7 2.94 ´ 10-6 7.35 ´ 10-7
16. For the reaction A + 2B ® C, the rate law is
. What are the units of the rate constant where time is measured in seconds? A. B. C.
D. E.
17. The reaction of NO and O2 produces NO2. 2 NO(g) + O2(g) ® 2 NO2(g) The reaction is second-order with respect to NO(g) and first-order with respect to O2(g). At a given temperature, the rate constant, k, equals 4.7 ´ 102 M–2s–1. What is the rate of reaction when the initial concentrations of NO and O2 are 0.025 M and 0.015 M, respectively? A. 2.6 ´ 10–3 M/s B. 4.4 ´ 10–3 M/s C. 0.18 M/s D. 2.0 ´ 10–8 M/s E. 3.8 ´ 102 M/s
18. For the second-order decomposition reaction, 2A ® B which of the following relationships yields a straight line plot?
A. ln [A] versus k B. ln [A]2 versus t C. ln [A] versus t D. 1/[A] versus t E. 1/[A] versus k
rate = k[A]2
19. Which of the following expressions corresponds to the integrated rate equation for a first-order decomposition reaction? A. [A] = -kt B. [A] = [A]0 -kt C. ln [A] = ln [A]0 -kt D. E.
20. The rate constant of a first-order decomposition reaction is 0.0147 s–1. If the initial concentration of reactant is 0.178 M, what is the concentration of reactant after 30.0 seconds? A. 8.72 ´ 105 M B. 0.0645 M C. 0.115 M D. 0.0785 M E. 0.643 M
21. The rate constant for the decomposition of cyclobutane is 2.08 ´ 10-2 s-1 at high temperatures. C4H8(g) ® 2C2H4(g) How many seconds are required for an initial concentration of 0.100 M C4H8(g) to decrease to 0.0450 M? A. 0.00114 s B. 1.07 s C. 2.64 s D. 38.4 s E. 874 s
22. At a high temperature, the first-order decomposition of N2O5(g) produces NO2(g) and O2(g). If the initial concentration of 0.400 M N2O5(g) is reduced to 0.169 M after 118 seconds, what is the rate constant for the reaction? A. 1.96 ´ 10-3 s-1 B. 2.29 ´ 10-3 s-1 C. 4.37 ´ 10-3 s-1 D. 7.30 ´ 10-3 s-1 E. 1.37 ´ 102 s-1
23. The reaction A ® B follows first-order kinetics with a half-life of 21.7 hours. If the concentration of A is 0.023 M after 48.0 hours, what is the initial concentration of A? A. 0.0050 M B. 0.051 M C. 0.51 M D. 0.11 M E. 2.0 ´ 102 M
24. For the first-order reaction below, the concentration of product B after 24.2 seconds is 0.322 M. If k = 8.75 ´ 10-2 s-1, what was the initial concentration of A? A ® 2B
rate = k[A]
A. 0.0341 M B. 0.183 M C. 1.34 M D. 2.68 M E. 29.3 M 25. Which equation is used to calculate the half-life of the first-order equation below? A ® 2B
rate = k[A]
A. t1/2 = 2 ´ ln k B. t1/2 = (ln k)/[A]o C. t1/2 = ln (k/2) D. t1/2 = ln ([A]o/k) E. t1/2 = (ln 2)/k 26. Which equation is used to calculate the half-life of the second-order equation below? 2A ® B
rate = k[A]2
A.
B.
C.
D. E. 27. Which equation is used to calculate the half-life of the zero-order equation below? A ® 2B
rate = k
A.
B.
C. D. E. t1/2 = k 28. For the second-order reaction below, the rate constant of the reaction is 9.4 ´ 10–3 M–1s–1. How long (in seconds) is required to decrease the concentration of A from 2.16 M to 0.40 M? 2A ® B
A. 2.0 ´ 101 s B. 7.8 ´ 101 s C. 1.8 ´ 102 s D. 1.9 ´ 102 s E. 2.2 ´ 102 s
rate = k[A]2
29. For the second-order reaction below, the initial concentration of A is 1.34 M. If k = 7.61 ´ 10-4 s-1, what is the concentration of A after 18.3 minutes? 2A ® B
rate = k[A]2
A. 0.504 M B. 0.581 M C. 0.632 M D. 0.836 M E. 1.12 M 30. For the second-order reaction below, the initial concentration of A is 1.00 M. If the concentration of A is reduced to 0.43 M after 75 seconds, what is the rate constant? 2A ® B
rate = k[A]2
A. 5.2 ´ 10-4 L/mols B. 5.7 ´ 10-3 L/mols C. 7.6 ´ 10-3 L/mols D. 1.1 ´ 10-2 L/mols E. 1.8 ´ 10-2 L/mols 31. For the second-order reaction below, the concentration of product B after 132 seconds is 0.0281 M. If the initial concentration of A is 0.932 M, what is the rate constant? 2A ® B
rate = k[A]2
A. 1.43 ´ 10-5 L/mols B. 4.71 ´ 10-4 L/mols C. 5.22 ´ 10-4 L/mols D. 6.63 ´ 10-3 L/mols E. 96.4 L/mols 32. What is the half-life of a first-order reaction if the rate constant is 6.2 ´ 10–3 s–1? A. 8.9 ´ 10–3 s B. 0.097 s C. 5.77 s D. 1.0 ´ 102 s E. 1.6 ´ 102 s
33. For the first-order reaction below, the initial concentration of A is 0.240 M. If the concentration of A decreases to 0.0800 M after 21.8 hours, what is the half-life of the reaction? A®B
rate = k[A]
A. 0.0504 hrs B. 1.28 hrs C. 7.26 hrs D. 10.4 hrs E. 13.8 hrs 34. Hydrogen peroxide decomposes into water and oxygen in a first-order process. H2O2(aq) ® H2O(l) + 1/2 O2(g) At 20.0C, the half-life for the reaction is 3.92 ´ 104 seconds. If the initial concentration of hydrogen peroxide is 0.52 M, what is the concentration after 7.00 days? A. 1.2 ´ 10-5 M B. 0.034 M C. 0.074 M D. 0.22 M E. 0.52 M
35. A first-order reaction has a half-life of 4.54 seconds. How much time is required for the reactant to be reduced to 6.25% of its initial concentration? A. 1.14 s B. 9.08 s C. 13.6 s D. 18.2 s E. 93.6 s
36. For the first-order reaction below, the initial concentration of A is 0.80 M. What is the half-life of the reaction if the concentration of A decreases to 0.10 M in 54 seconds? A®B
A. 18 s B. 24 s C. 36 s D. 48 s E. 51 s
rate = k[A]
37. A first-order reaction has a half-life of 2.10 hours. How much time is required for the reactant to be reduced to 33.0% of its initial concentration? A. 2.45 hours B. 2.89 hours C. 3.36 hours D. 3.93 hours E. 4.21 hours
38. The decomposition of formic acid follows first-order kinetics. HCO2H(g) ® CO2(g) + H2(g) The half-life for the reaction at 550C is 24 seconds. How many seconds does it take for the formic acid concentration to decrease by 87.5%? A. 24 s B. 36 s C. 48 s D. 72 s E. 96 s
39. Which of the following factors are likely to affect the rate of a chemical reaction? 1. 2. 3.
the presence of a catalyst the temperature of the reactants the physical state (solid, liquid, or gas) of the reactants
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 40. The effect of a catalyst on a chemical reaction is to A. enable an alternate path for the reaction that has a lower activation barrier. B. increase the energy of the products. C. increase the activation barrier for the forward reaction. D. increase the energy of the reactants. E. increase the frequency of collisions between reactants and products.
41. The transition-state model for reaction rate assumes that the activated complex 1. 2. 3.
is in equilibrium with the reactants. is gaseous at low concentrations. may decompose to form products or may revert to reactants.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 42. The correct form of the Arrhenius equation is A. Ea = Ae-k/RT B. C. D. Ea = Aek/RT E. A = Eae-k/RT
43. For a given reaction, the rate constant doubles when the temperature is increased from 45.0C to 73.0C. What is the activation energy for this reaction? (R = 8.31 J/Kmol) A. 0.676 kJ/mol B. 9.85 kJ/mol C. 16.1 kJ/mol D. 22.7 kJ/mol E. 65.4 kJ/mol
44. For a given reaction, the activation energy is 63.9 kJ/mol. If the reaction rate constant is 4.1 ´ 10–3 L/mols at 32.0C, what is the reaction rate constant at –5.0C? (R = 8.31 J/Kmol) A. 1.2 ´ 10–3 L/mols B. 1.4 ´ 10–6 L/mols C. 1.3 ´ 10–4 L/mols D. 4.0 ´ 10–3 L/mols E. 1.3 ´ 10–1 L/mols
45. The rate constant at 373 K for a certain reaction is 8.29 ´ 10-4 s-1 and the activation energy is 12.0 kJ/mole. What is the value of the constant, A, in the Arrhenius equation? A. 0.00832 s-1 B. 0.0398 s-1 C. 0.120 s-1 D. 0.998 s-1 E. 25.1 s-1
46. For a certain reaction, the activation energy is 52.1 kJ/mole. By what ratio will the rate constant change if the temperature is decreased from 175C to 75C? A. 0.00402 B. 0.0179 C. 0.996 D. 121 E. 249
47. All of the following statements are correct EXCEPT A. a heterogeneous catalyst is in a different phase from the reaction mixture. B. enzymes are protein molecules that catalyze reactions. C. catalysts are not consumed in reactions. D. a reaction mechanism describes the path of a reaction at the molecular level. E. elementary steps in a reaction mechanism are always unimolecular.
48. The elementary steps for the catalyzed decomposition of dinitrogen monoxide are shown below. N2O(g) + NO(g) ® N2(g) + NO2(g) NO2(g) ® NO(g) + 1/2O2(g) Which of the following statements is/are CORRECT? 1. 2. 3.
The overall balanced reaction is N2O(g) ® N2(g) + 1/2O2(g). NO2(g) is a catalyst for the reaction. NO(g) is a reaction intermediate.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 49. The elementary steps for the catalyzed decomposition for dinitrogen monoxide are shown below. Identify the catalyst. NO(g) + N2O(g) ® N2(g) + NO2(g) 2NO2(g) ® 2NO(g) + O2(g) A. NO(g) B. N2O(g) C. N2(g) D. NO2(g) E. O2(g)
50. The elementary steps for a catalyzed reaction are shown below. Identify the catalyst. Identify the reactive intermediate. H2O2(aq) + I-(aq) ® H2O(l) + IO-(aq) IO-(aq) + H2O2(aq) ® H2O(l) + O2(g) + I-(aq) A. The catalyst is H2O(l); the reactive intermediate is I-(aq). B. The catalyst is IO-(aq); the reactive intermediate is I-(aq). C. The catalyst is I-(aq); the reactive intermediate is H2O2(aq). D. The catalyst is I-(aq); the reactive intermediate is IO-(aq). E. The catalyst is H2O2(aq); the reactive intermediate is I-(aq).
51. A possible mechanism for the decomposition of ozone to oxygen in the atmosphere is O3(g) O2(g) + O O + O3(g) ® 2O2(g)
(fast equilibrium) (slow)
What is a rate law that is consistent with this mechanism?
A. rate = k[O3] B. rate = k[O3]2 C. rate = k[O3] ´ [O] D. rate = k[O3]2 ´ [O2] E. rate = k[O3]2 ´ [O2]-1 52. A possible reaction mechanism for the reaction of nitrogen dioxide with carbon monoxide is 2NO2(g) ® NO3(g) + NO(g) NO 3(g) + CO(g) ® NO2(g) + CO2(g)
(slow) (fast)
What is the overall reaction and the most probable rate law for the reaction?
A. NO2(g) + CO(g) ® NO(g) + CO2(g); rate = k[NO2] B. NO2(g) + CO(g) ® NO(g) + CO2(g); rate = k[NO2]2 C. NO2(g) + CO(g) ® NO(g) + CO2(g); rate = k[NO3]´[CO] D. 2NO2(g) + CO(g) ® NO(g) + CO2(g); rate = k[NO2]´[ NO3]´[CO] E. 2NO2(g) + CO(g) ® NO(g) + CO2(g); rate = k[NO3]´[CO] 53. Nitrogen dioxide reacts with carbon monoxide to produce nitrogen monoxide and carbon dioxide. NO2(g) + CO(g) ® NO(g) + CO2(g) A proposed mechanism for this reaction is 2NO2(g) NO3(g) + NO(g) NO3(g) + CO(g) ® NO2(g) + CO2(g) What is a rate law that is consistent with the proposed mechanism?
A. rate = k[NO2]2´[CO]´ [NO]–1 B. rate = k[NO2]2´[CO] C. rate = k[NO2]´[CO] D. rate = k[NO3]´[CO] E. rate = k[NO2]2
(fast, equilibrium) (slow)
54. For the overall reaction A + 2B ® C which of the following mechanisms yields the correct overall chemical equation and is consistent with the rate equation below? rate = k[A]´[B] A. A + B I (fast) I+A®C (slow) B. A + B ® I (slow) I+B®C (fast) C. 2B ® I (slow) A+I®C (fast) D. 2B I (fast) I+A®C (slow) E. A + 2B I (fast) I+B®C+B (slow)
55. For the overall reaction 2A + B ® C which of the following mechanisms are consistent with a rate equation of rate = k[A]2´[B]? A. A + B I (fast) I + A ® C (slow) B. A + B ® I (slow) I + A ® C (fast) C. 2A ® I (slow) B + I ® C (fast) D. 2A I (fast) I + B ® C (slow) E. Answers a and d are both correct.
56. What did Henry Eyring develop that changed how scientists study rates of reactions? A. the catalyst model B. the transition-state model C. the liquid-state model D. the temperature dependence model E. the enzyme model
57. Ozone in the upper atmosphere is important to life on earth because it converts harmful UV radiation to heat. How does the release of CFCs into the atmosphere affect the concentration of ozone in the upper atmosphere? A. the Cl in the CFC catalyzes the decomposition of ozone decreasing its concentration B. the Cl in the CFC catalyzes the production of ozone increasing its concentration C. the NO2 in the CFC catalyzes the decomposition of ozone decreasing its concentration D. the NO2 in the CFC catalyzes the production of ozone increasing its concentration E. the CFC's change the ozone into photochemical smog
Chapter 12--Gaseous Chemical Equilibrium 1. All of the following statements are false for a chemical system in a dynamic equilibrium EXCEPT A. the concentrations of reactants and products must be equal. B. the forward reaction is endothermic. C. the forward reaction is exothermic. D. the chemical reaction proceeds in the forward direction until all reactants are consumed. E. the concentrations of reactants and products remain constant over time.
2. The partial pressure of a gas is A. directly proportional to the number of moles of the gas. B. proportional to the gas constant, R. C. directly proportional to the volume of the gas. D. always constant during a chemical reaction. E. inversely proportional to the temperature of the gas.
3. Which of the following statements is/are CORRECT? 1. 2. 3.
Product concentrations appear in the numerator of an equilibrium constant expression. A reaction favors the formation of products if K >> 1. Stoichiometric coefficients are used as exponents in an equilibrium constant expression.
A. 1 only B. 2 only C. 3 only D. 2 and 3 E. 1, 2, and 3
4. What is the correct equilibrium constant expression for the formation of ammonia gas from nitrogen gas and hydrogen gas?
A.
B.
C.
D.
E.
5. What is the correct equilibrium constant expression for the following reaction? CO2(g) + 2H2O(g) CH4(g) + 2O2(g)
A.
B.
C.
D. E. none of the above
6. What is the correct equilibrium constant expression for the following reaction? C4H10(g) +
A.
B.
C.
D.
E.
O2(g)
4CO2(g) + 5H2O(g)
7. What is the correct equilibrium constant expression for the following reaction? H2(g) + I2(s) 2HI(g)
A.
B.
C.
D.
E.
8. What is the correct equilibrium constant expression for the following reaction? MgCO3(s) MgO(s) + CO2(g)
A.
B. C. D. E.
9. Write a balanced chemical equation which corresponds to the following equilibrium constant expression.
A. 1/2N2(g) + 3/2H2(g) NH3(g) B. N2(g) + 3 H2(g) 2NH3(g) C. 2NH3(g) N2(g) + 3H2(g) D. NH3(g) 1/2N2(g) + 3/2H2(g) E. 2N2(g) + 6H2(g) 4NH3(g)
10. Write the balanced chemical reaction which corresponds to the following equilibrium constant expression.
A. 4N2O5(g) O2(g) + 2NO2(g) B. 2N2O5(g) O2(g) + 4NO2(g) C. N2O5(g) O2(g) + NO2(g) D. O2(g) + 2NO2(g) 4N2O5(g) E. O2(g) + 4NO2(g) 2N2O5(g)
11. The value of the equilibrium constant for the following reaction is 345. A + 2B 3C + D What is the value of the equilibrium constant for the reaction below? 2A + 4B 6C + 2D A. K = 345 B. K = (345)2 = 1.19 ´ 105 C. K = (345)1/2 = 18.6 D. K = (2 ´ 345)2 = 4.76 ´ 105 E. K = 2 ´ (345)2 = 2.38 ´ 105
12. For the following reaction, 2SO3(g) 2SO2(g) + O2(g) the equilibrium constant, K, is 1.32 at 627C. What is the equilibrium constant, at 627C, for the reaction below? SO2(g) + 1/2O2(g) SO3(g) A. –1.15 B. –0.66 C. 0.379 D. 0.870 E. 1.52
13. Given the following chemical equilibria, N2(g) + O2(g) N2(g) + 3 H2(g) H2(g) + 1/2O2(g)
2NO(g) 2NH3(g) H2O(g)
Determine the method used to calculate the equilibrium constant for the reaction below. 4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g)
K1 K2 K3
Kc
A.
B. C.
D.
E. 14. Use the equilibrium constants for the following reactions 2NO(g) N2(g) + O2(g) 2NO(g) + O2(g) 2NO2(g)
K1 = 2.4 ´ 1030 K2 = 2.4 ´ 1012
to determine the equilibrium constant for the reaction below. N2(g) + 2O2(g) 2NO2(g)
A. 1.7 ´ 10-43 B. 1.0 ´ 10-18 C. 5.8 ´ 1018 D. 2.4 ´ 1030 E. 5.8 ´ 1042 15. Use the equilibrium constants for the following reactions at 700C 2SO2(g) + O2(g) 2NO(g) + O2(g)
2SO3(g) 2NO2(g)
K1 = 4.8 K2 = 16
to determine the equilibrium constant for the following reaction. SO3(g) + NO(g) SO2(g) + NO2(g)
A. 0.30 B. 0.55 C. 0.85 D. 1.8 E. 3.3 16. Given the following equilibrium equations, 2N2O(g) N2O4(g) 2NO2(g)
2N2(g) + O2(g) 2NO2(g) N2(g) + 2O2(g)
calculate K for the decomposition of dinitrogen tetraoxide to nitrogen dioxide and oxygen. 2N2O4(g) 2N2O(g) + 3O2(g)
A. 1.1 ´ 10-41 B. 1.3 ´ 10-20 C. 8.9 ´ 10-7 D. 3.3 ´ 1021 E. 2.3 ´ 1049 17. For which of the following reactions does Kc equal Kp? A. Sn(s) + 2H2O(g) SnO2(s) + 2H2(g) B. 2C2H6(g) + 7O2(g) 4CO2(g) + 6H2O(g) C. NH4Cl(s) NH3(g) + HCl(g) D. N2(g) + 3H2(g) 2NH3(g) E. CaCO3(s) CaO(s) + CO2(g)
K1 = 8.3 ´ 1034 K2 = 4.6 ´ 10-3 K3 = 5.9 ´ 1016
K4 = ?
18. What is the relationship between Kp and Kc for the reaction below? N2(g) + 3 H2(g) 2 NH3(g)
A.
B. C.
D.
E.
19. Which of the following reactions is a heterogeneous equilibrium expression? A. 2NO(g) + O2(g) 2NO2(g) B. 2NH3(g) N2(g) + 3H2(g) C. 2H2(g) + O2(g) 2H2O(g) D. 2S(s) + 3O2(g) 2SO3(g) E. C2H4(g) + H2(g) C2H6(g)
20. Which of the following reactions is a homogeneous equilibrium expression? A. CaCO3(s) CaO(s) + CO2(g) B. NH3(g) + HCl(g) NH4Cl(s) C. Mg(s) + Cl2(g) MgCl2(s) D. FeO(s) + CO(g) Fe(s) + CO2(g) E. 2NO(g) + O2(g) 2NO2(g)
21. The reaction below was studied at a high temperature. At equilibrium, the partial pressures of the gases are as follows: PCl5 = 1.4 ´ 10-4 atm, PCl3 = 2.4 ´ 10-2 atm, Cl2 = 3.0 ´ 10-1 atm. What is the value of K for the reaction? PCl5(g) PCl3(g) + Cl2(g) A. 4.3 ´ 10-4 B. 0.019 C. 0.32 D. 51 E. 2.3 ´ 103
22. For the reaction below, the partial pressures of gases at equilibrium are as follows: H2 = 7.1 ´ 10-5 atm, Cl2 = 2.5 ´ 10-6 atm, and HCl = 3.0 atm. What is the value of the equilibrium constant, K? H2(g) + Cl2(g) 2HCl(g) A. 2.0 ´ 10-11 B. 5.9 ´ 10-11 C. 1.6 ´ 10-9 D. 1.7 ´ 1010 E. 5.1 ´ 1010
23. At 25C, the partial pressure of gases at equilibrium are as follows: N2 = 0.12 atm, O2 = 0.040 atm, and NO = 4.5 ´ 10-17 atm. What is the value of the equilibrium constant, K? N2(g) + O2(g) 2NO(g) -31 A. 4.2 ´ 10 B. 5.2 ´ 10-21 C. 9.4 ´ 10-15 D. 1.1 ´ 1014 E. 2.4 ´ 1030
24. Nitrogen can react with oxygen to form nitrogen monoxide. N2(g) + O2(g)
2NO(g)
K = 1.0 ´ 10-30 at 25C
When equilibrium is established, the partial pressures of nitrogen and oxygen are 1.2 atm and 3.1 atm, respectively. What is the equilibrium pressure of nitrogen monoxide?
A. 3.7 ´ 10-30 atm B. 1.9 ´ 10-15 atm C. 2.7 ´ 10-13 atm D. 1.9 ´ 10-10 atm E. 3.7 ´ 10-7 atm
25. At 25C, the decomposition of dinitrogen tetraoxide N2O4(g) 2NO2(g) has an equilibrium constant (K) of 0.144. If the equilibrium pressure of nitrogen dioxide is 0.298 atm, what is the pressure of dinitrogen tetraoxide? A. 0.0128 atm B. 0.617 atm C. 1.03 atm D. 1.62 atm E. 2.07 atm
26. For the following reaction, the equilibrium constant (K) equals 21.2. SnO2(s) + 2H2(g) Sn(s) + 2H2O(g) At equilibrium, the total pressure of the system is 0.390 atm. What is the partial pressure of each gas? A. 0.0303 atm H2; 0.360 atm H2O B. 0.0522 atm H2; 0.339 atm H2O C. 0.0696 atm H2; 0.320 atm H2O D. 0.320 atm H2; 0.0696 atm H2O E. 0.339 atm H2; 0.0522 atm H2O
27. At 25C, the decomposition of dinitrogen tetraoxide N2O4(g) 2NO2(g) has an equilibrium constant (K) of 0.144. At equilibrium, the total pressure of the system is 0.500 atm. What is the partial pressure of each gas? A. N2O4 = 0.206 atm; NO2 = 0.294 atm B. N2O4 = 0.212 atm; NO2 = 0.288 atm C. N2O4 = 0.288 atm; NO2 = 0.212 atm D. N2O4 = 0.294 atm; NO2 = 0.206 atm E. N2O4 = 0.437 atm; NO2 = 0.063 atm
28. At sufficiently high temperatures, ammonium iodide decomposes to ammonia and hydrogen iodide. NH4I(s) NH3(g) + HI(g) A mass of 5.00 g of NH4I is sealed in a 2.00-L flask and heated to 673 K. If 2.56 g NH4I(s) remain unreacted when the system has reach equilibrium, what is the equilibrium constant (Kp) for the reaction? (R = 0.0821 Latm/molK) A. 0.15 B. 0.22 C. 0.47 D. 0.59 E. 0.89
29. Hydrogen iodide can decompose into hydrogen and iodine gases. 2HI(g) H2(g) + I2(g) K for the reaction is 0.016. If 0.148 atm of HI(g) is sealed in a flask, what is the pressure of each gas when equilibrium is established? A. HI = 0.118 atm; H2 = 0.015 atm; I2 = 0.015 atm B. HI = 0.133 atm; H2 = 0.015 atm; I2 = 0.015 atm C. HI = 0.110 atm; H2 = 0.019 atm; I2 = 0.019 atm D. HI = 0.126 atm; H2 = 0.022 atm; I2 = 0.022 atm E. HI = 0.174 atm; H2 = 0.022 atm; I2 = 0.022 atm
30. Hydrogen iodide can decompose into hydrogen and iodine gases. 2HI(g) H2(g) + I2(g) Kp for the reaction is 0.016. If 0.350 atm of HI(g) is sealed in a flask, what is the total pressure of the system when equilibrium is established? A. 0.258 atm B. 0.279 atm C. 0.350 atm D. 0.385 atm E. 0.412 atm
31. Nitrosyl bromide decomposes according to the chemical equation below. 2NOBr(g) 2NO(g) + Br2(g) 1.00 atm of NOBr is sealed in a flask. At equilibrium, the partial pressure of NOBr is 0.82 atm. What is the equilibrium constant for the reaction? A. 3.6 ´ 10-3 B. 8.7 ´ 10-3 C. 2.8 ´ 10-2 D. 3.5 ´ 10-2 E. 4.3 ´ 10-3
32. Nitrosyl chloride decomposes according to the chemical equation below. 2NOCl(g) 2NO(g) + Cl2(g) A pressure of 0.320 atm of nitrosyl chloride is sealed in a flask and allowed to reach equilibrium. If 22.6% of the NOCl decomposes, what is the equilibrium constant for the reaction? A. 0.00153 B. 0.00308 C. 0.00611 D. 0.00730 E. 0.02471
33. At a given temperature, the equilibrium constant (Kp) for the decomposition of dinitrogen tetraoxide to nitrogen dioxide is 0.172. If 0.224 atm N2O4 is sealed in a flask, what partial pressure of NO2 will exist at equilibrium? N2O4(g) 2NO2(g) A. 0.0385 atm B. 0.158 atm C. 0.196 atm D. 0.257 atm E. 0.379 atm
34. The reaction of nitrogen gas and oxygen gas to form nitrogen monoxide, N2(g) + O2(g) 2NO(g) has an equilibrium constant of 1.0 ´ 10-30 at 298 K. What equilibrium partial pressure of NO(g) will form if 0.50 atm of N2 and 0.50 atm of O2 are sealed in a flask at 298 K? A. 1.0 ´ 10-60 atm B. 5.0 ´ 10-31 atm C. 1.0 ´ 10-30 atm D. 1.0 ´ 10-15 atm E. 5.0 ´ 10-16 atm
35. For the system CO(g) + H2O(g) CO2(g) + H2(g) K is 1.6 at 900 K. If 0.400 atm CO(g) and 0.400 atm H2O(g) are combined in a sealed flask, what is the equilibrium partial pressure of CO2(g)? A. 0.22 atm B. 0.31 atm C. 0.47 atm D. 0.51 atm E. 0.65 atm
36. Consider the following equilibrium: N2(g) + O2(g) 2NO(g) At a certain temperature the equilibrium constant for the reaction is 0.0255. What is the partial pressure of NO gas at equilibrium if the initial pressure of all the gases (both reactants and products) is 0.300 atm? A. 6.65 ´ 10-2 B. 0.183 C. 0.234 D. 0.252 E. 0.417
37. For the reaction 2A 3B Kc = 1.37. If the concentrations of A and B are equal, what is the value of that concentration? A. 0.685 M B. 0.822 M C. 1.17 M D. 1.37 M E. 1.88 M
38. Which of the following may change the ratio of products to reactants in an equilibrium mixture for a chemical reaction involving gaseous species? 1. 2. 3.
Increasing the temperature. Adding a catalyst. Adding gaseous reactants.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1 and 3 39. Assume that the following chemical reaction is at equilibrium. 2ICl(g)
I2(g) + Cl2(g)
DH = +26.9 kJ
At 25C, K = 2.0 ´ 105. If the temperature is increase to 45C, which statement applies?
A. K will decrease and the reaction will proceed in the backward direction. B. K will decrease and the reaction will proceed in the forward direction. C. K will remain unchanged and the reaction will proceed in the forward direction. D. K will increase and the reaction will proceed in the backward direction. E. K will increase and the reaction will proceed in the forward direction. 40. Assuming the reaction below is at equilibrium, which of the following changes will drive the reaction to the left? C(s) + O2(g)
1. 2. 3.
CO2(g)
Increasing the temperature. Adding O2(g). Removing C(s).
DH = -393.5 kJ/mol
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 2 and 3 41. The formation of ammonia from elemental nitrogen and hydrogen is an exothermic process. N2(g) + 3H2(g)
2NH3(g)
DH = -92.2 kJ/mol
Assuming the reaction is at equilibrium, which one of the following changes will drive the reaction to the right?
A. adding ammonia B. increasing the temperature C. increasing the pressure D. removing hydrogen E. adding a catalyst 42. Assume that the following endothermic chemical reaction is at equilibrium. C(s) + H2O(g) H2(g) + CO(g) Which of the following statements is/are CORRECT? 1. 2. 3.
Increasing the concentration of H2(g) will cause the reaction to proceed in the backward direction, increasing the equilibrium concentration of H2O(g). Decreasing the temperature will cause the reaction to proceed in the forward direction, increasing the equilibrium concentration of CO(g). Increasing the amount of C(s) will cause the reaction to proceed in the forward direction, increasing the equilibrium concentration of CO(g).
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 43. In which of the following equilibrium systems will an increase in the pressure have no effect on the concentrations of products and reactants? A. H2(g) + F2(g) 2 HF(g) B. N2(g) + 3 H2(g) 2 NH3(g) C. CaCO3(s) CaO(s) + CO2(g) D. 2 NOBr(g) 2 NO(g) + Br2(g) E. 2 H2O(g) + O2(g) 2 H2O2(g)
44. In which of the following equilibrium systems would an increase in volume (at constant temperature) cause the reaction to shift to the right? A. N2O4(g) 2NO2(g) B. N2(g) + 3H2(g) 2NH3(g) C. H2(g) + Cl2(g) 2HCl(g) D. Answers a and b are correct. E. Answers b and c are correct.
45. The equilibrium constant for a gas phase reaction is measured at two temperatures. At 100C, the equilibrium constant is 36. At 200C, the equilibrium constant is 147. Which of the following statements is correct for this equilibrium? A. The reaction must be first-order. B. The reaction is endothermic. C. A catalyst must be present. D. Each reactant molecule decomposes into two or more product molecules. E. One of the products must be a solid.
46. Which of the following statements is/are CORRECT for the following system at equilibrium? H2(g) + F2(g) 2 HF(g) 1. 2. 3.
Addition of a catalyst will increase the equilibrium constant, causing the reaction to proceed in the forward direction. The values of Kc and Kp are identical. Increasing the pressure of a system will cause the reaction to proceed in the forward direction.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 47. The equilibrium constant, Kp, for the reaction below is 0.24 at 1500C. SO3(g) + NO(g) SO2(g) + NO2(g) If 0.30 atm of sulfur trioxide, 0.15 atm of nitrogen monoxide, 0.55 atm of sulfur dioxide, and 0.030 atm of nitrogen dioxide are mixed, what changes in pressure will occur? A. The pressures of SO3 and NO decrease; the pressures of SO2 and NO2 increase. B. The pressures of SO3 and NO increase; the pressures of SO2 and NO2 decrease. C. The pressures of SO3 and SO2 decrease; the pressures of NO and NO2 increase. D. The pressures of SO3 and SO2 increase; the pressures of NO and NO2 decrease. E. Equal numbers of particles exist on both sides of the equation; no reaction will occur.
48. If the value of Q is less than Kp, then A. the system is in equilibrium. B. a catalyst is necessary to achieve equilibrium. C. the reaction will go left or right depending upon the reaction stoichiometry. D. the reaction will proceed to the right until equilibrium is established. E. the reaction will proceed to the left until equilibrium is established.
49. A 2.5 L flask is filled with 0.25 atm SO3, 0.20 atm SO2, and 0.40 atm O2, and allowed to reach equilibrium. Assume the temperature of the mixture is chosen so that Kp = 0.12. Predict the effect on the partial pressure of SO3 as equilibrium is achieved by using Q, the reaction quotient. 2 SO3(g) 2 SO2(g) + O2(g) A. The partial pressure of SO3 will decrease because Q > K. B. The partial pressure of SO3 will decrease because Q < K. C. The partial pressure of SO3 will increase because Q < K. D. The partial pressure of SO3 will increase because Q > K. E. The partial pressure of SO3 will remain the same because Q = K.
50. For the reaction C(s) + CO2(g) 2CO(g) Kc = 168. A mixture contains some C(s), [CO] = 0.50 M and [CO2] = 0.75 M. Therefore the system ____ at equilibrium, because ____. A. is not; the value of Q is 0.67 B. is not; the value of Q is 1.5 C. is; the value of Q is 0.67 D. is not; the value of Q is 0.33 E. is; the value of Q is 0.33
51. Consider the reaction A(g) 2B(g) where Kp = 4.1 at 25C. If 0.75 atm A(g) and 1.5 atm B(g) are initially present in a 1.0 L flask at 25C, what change in partial pressures (if any) will occur in time? A. The partial pressure of A will decrease and the partial pressure of B will decrease. B. The partial pressure of A will decrease and the partial pressure of B will increase. C. The partial pressure of A will increase and the partial pressure of B will decrease. D. The partial pressure of A will increase and the partial pressure of B will increase. E. The partial pressures of both A and B will remain unchanged.
52. "If a chemical system at equilibrium is disturbed by adding a gaseous species (reactant or product), the reaction will proceed in such a direction as to consume part of the added species" is a statement of A. the ideal gas law. B. Le Châtelier's principle. C. the de Broglie equation. D. the van't Hoff equation. E. the first law of thermodynamics.
53. Calcium carbonate decomposes to calcium oxide and carbon dioxide. CaCO3(s)
CaO(s) + CO2(g)
DH = +179 kJ
The equilibrium constant for this reaction is 9.7 ´ 10-24 at 298 K. What is the equilibrium constant at 575 K? (R = 8.31 J/molK)
A. 7.5 ´ 10-16 B. 1.3 ´ 10-8 C. 1.4 ´ 1038 D. 1.3 ´ 1015 E. 1.0 ´ 1023 54. The reaction of nitrogen with hydrogen to form ammonia is thermodynamically favorable. N2(g) + 3H2(g)
2NH3(g)
DH = -92.2 kJ
The equilibrium constant for this reaction is 6.0 ´ 105 at 298 K. At what temperature is the equilibrium constant equal to 1.0 ´ 103? (R = 8.31 J/molK)
A. 85 K B. 110 K C. 310 K D. 360 K E. 2800 K 55. The Haber process for the production of ammonia relies on high temperatures and pressures. Which of these, high temperatures or pressures, actually reduce the yield of the reaction at equilibrium? N2(g) + 3H2(g)
A. high pressure B. high temperature C. both D. neither E. can't be determined
2NH3(g)
DH = -92.2 kJ
56. The Haber process for the production of ammonia relies on a heterogeneous catalyst. How does the use of this catalyst effect the yield of the reaction at equilibrium? N2(g) + 3H2(g)
2NH3(g)
A. yield is increased B. yield is decreased C. depends on the catalyst used D. doesn't effect yield E. can't be determined
DH = -92.2 kJ
Chapter 12--Gaseous Chemical Equilibrium Key
1. All of the following statements are false for a chemical system in a dynamic equilibrium EXCEPT A. the concentrations of reactants and products must be equal. B. the forward reaction is endothermic. C. the forward reaction is exothermic. D. the chemical reaction proceeds in the forward direction until all reactants are consumed. E. the concentrations of reactants and products remain constant over time.
2. The partial pressure of a gas is A. directly proportional to the number of moles of the gas. B. proportional to the gas constant, R. C. directly proportional to the volume of the gas. D. always constant during a chemical reaction. E. inversely proportional to the temperature of the gas.
3. Which of the following statements is/are CORRECT? 1. 2. 3.
Product concentrations appear in the numerator of an equilibrium constant expression. A reaction favors the formation of products if K >> 1. Stoichiometric coefficients are used as exponents in an equilibrium constant expression.
A. 1 only B. 2 only C. 3 only D. 2 and 3 E. 1, 2, and 3
4. What is the correct equilibrium constant expression for the formation of ammonia gas from nitrogen gas and hydrogen gas?
A.
B.
C.
D.
E.
5. What is the correct equilibrium constant expression for the following reaction? CO2(g) + 2H2O(g) CH4(g) + 2O2(g)
A.
B.
C.
D. E. none of the above
6. What is the correct equilibrium constant expression for the following reaction? C4H10(g) +
A.
B.
C.
D.
E.
O2(g)
4CO2(g) + 5H2O(g)
7. What is the correct equilibrium constant expression for the following reaction? H2(g) + I2(s) 2HI(g)
A.
B.
C.
D.
E.
8. What is the correct equilibrium constant expression for the following reaction? MgCO3(s) MgO(s) + CO2(g)
A.
B. C. D. E.
9. Write a balanced chemical equation which corresponds to the following equilibrium constant expression.
A. 1/2N2(g) + 3/2H2(g) NH3(g) B. N2(g) + 3 H2(g) 2NH3(g) C. 2NH3(g) N2(g) + 3H2(g) D. NH3(g) 1/2N2(g) + 3/2H2(g) E. 2N2(g) + 6H2(g) 4NH3(g)
10. Write the balanced chemical reaction which corresponds to the following equilibrium constant expression.
A. 4N2O5(g) O2(g) + 2NO2(g) B. 2N2O5(g) O2(g) + 4NO2(g) C. N2O5(g) O2(g) + NO2(g) D. O2(g) + 2NO2(g) 4N2O5(g) E. O2(g) + 4NO2(g) 2N2O5(g)
11. The value of the equilibrium constant for the following reaction is 345. A + 2B 3C + D What is the value of the equilibrium constant for the reaction below? 2A + 4B 6C + 2D A. K = 345 B. K = (345)2 = 1.19 ´ 105 C. K = (345)1/2 = 18.6 D. K = (2 ´ 345)2 = 4.76 ´ 105 E. K = 2 ´ (345)2 = 2.38 ´ 105
12. For the following reaction, 2SO3(g) 2SO2(g) + O2(g) the equilibrium constant, K, is 1.32 at 627C. What is the equilibrium constant, at 627C, for the reaction below? SO2(g) + 1/2O2(g) SO3(g) A. –1.15 B. –0.66 C. 0.379 D. 0.870 E. 1.52
13. Given the following chemical equilibria, N2(g) + O2(g) N2(g) + 3 H2(g) H2(g) + 1/2O2(g)
2NO(g) 2NH3(g) H2O(g)
Determine the method used to calculate the equilibrium constant for the reaction below. 4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g)
K1 K2 K3
Kc
A.
B. C.
D.
E. 14. Use the equilibrium constants for the following reactions 2NO(g) N2(g) + O2(g) 2NO(g) + O2(g) 2NO2(g)
K1 = 2.4 ´ 1030 K2 = 2.4 ´ 1012
to determine the equilibrium constant for the reaction below. N2(g) + 2O2(g) 2NO2(g)
A. 1.7 ´ 10-43 B. 1.0 ´ 10-18 C. 5.8 ´ 1018 D. 2.4 ´ 1030 E. 5.8 ´ 1042 15. Use the equilibrium constants for the following reactions at 700C 2SO2(g) + O2(g) 2NO(g) + O2(g)
2SO3(g) 2NO2(g)
K1 = 4.8 K2 = 16
to determine the equilibrium constant for the following reaction. SO3(g) + NO(g) SO2(g) + NO2(g)
A. 0.30 B. 0.55 C. 0.85 D. 1.8 E. 3.3 16. Given the following equilibrium equations, 2N2O(g) N2O4(g) 2NO2(g)
2N2(g) + O2(g) 2NO2(g) N2(g) + 2O2(g)
calculate K for the decomposition of dinitrogen tetraoxide to nitrogen dioxide and oxygen. 2N2O4(g) 2N2O(g) + 3O2(g)
A. 1.1 ´ 10-41 B. 1.3 ´ 10-20 C. 8.9 ´ 10-7 D. 3.3 ´ 1021 E. 2.3 ´ 1049 17. For which of the following reactions does Kc equal Kp? A. Sn(s) + 2H2O(g) SnO2(s) + 2H2(g) B. 2C2H6(g) + 7O2(g) 4CO2(g) + 6H2O(g) C. NH4Cl(s) NH3(g) + HCl(g) D. N2(g) + 3H2(g) 2NH3(g) E. CaCO3(s) CaO(s) + CO2(g)
K1 = 8.3 ´ 1034 K2 = 4.6 ´ 10-3 K3 = 5.9 ´ 1016
K4 = ?
18. What is the relationship between Kp and Kc for the reaction below? N2(g) + 3 H2(g) 2 NH3(g)
A.
B. C.
D.
E.
19. Which of the following reactions is a heterogeneous equilibrium expression? A. 2NO(g) + O2(g) 2NO2(g) B. 2NH3(g) N2(g) + 3H2(g) C. 2H2(g) + O2(g) 2H2O(g) D. 2S(s) + 3O2(g) 2SO3(g) E. C2H4(g) + H2(g) C2H6(g)
20. Which of the following reactions is a homogeneous equilibrium expression? A. CaCO3(s) CaO(s) + CO2(g) B. NH3(g) + HCl(g) NH4Cl(s) C. Mg(s) + Cl2(g) MgCl2(s) D. FeO(s) + CO(g) Fe(s) + CO2(g) E. 2NO(g) + O2(g) 2NO2(g)
21. The reaction below was studied at a high temperature. At equilibrium, the partial pressures of the gases are as follows: PCl5 = 1.4 ´ 10-4 atm, PCl3 = 2.4 ´ 10-2 atm, Cl2 = 3.0 ´ 10-1 atm. What is the value of K for the reaction? PCl5(g) PCl3(g) + Cl2(g) A. 4.3 ´ 10-4 B. 0.019 C. 0.32 D. 51 E. 2.3 ´ 103
22. For the reaction below, the partial pressures of gases at equilibrium are as follows: H2 = 7.1 ´ 10-5 atm, Cl2 = 2.5 ´ 10-6 atm, and HCl = 3.0 atm. What is the value of the equilibrium constant, K? H2(g) + Cl2(g) 2HCl(g) A. 2.0 ´ 10-11 B. 5.9 ´ 10-11 C. 1.6 ´ 10-9 D. 1.7 ´ 1010 E. 5.1 ´ 1010
23. At 25C, the partial pressure of gases at equilibrium are as follows: N2 = 0.12 atm, O2 = 0.040 atm, and NO = 4.5 ´ 10-17 atm. What is the value of the equilibrium constant, K? N2(g) + O2(g) 2NO(g) -31 A. 4.2 ´ 10 B. 5.2 ´ 10-21 C. 9.4 ´ 10-15 D. 1.1 ´ 1014 E. 2.4 ´ 1030
24. Nitrogen can react with oxygen to form nitrogen monoxide. N2(g) + O2(g)
2NO(g)
K = 1.0 ´ 10-30 at 25C
When equilibrium is established, the partial pressures of nitrogen and oxygen are 1.2 atm and 3.1 atm, respectively. What is the equilibrium pressure of nitrogen monoxide?
A. 3.7 ´ 10-30 atm B. 1.9 ´ 10-15 atm C. 2.7 ´ 10-13 atm D. 1.9 ´ 10-10 atm E. 3.7 ´ 10-7 atm
25. At 25C, the decomposition of dinitrogen tetraoxide N2O4(g) 2NO2(g) has an equilibrium constant (K) of 0.144. If the equilibrium pressure of nitrogen dioxide is 0.298 atm, what is the pressure of dinitrogen tetraoxide? A. 0.0128 atm B. 0.617 atm C. 1.03 atm D. 1.62 atm E. 2.07 atm
26. For the following reaction, the equilibrium constant (K) equals 21.2. SnO2(s) + 2H2(g) Sn(s) + 2H2O(g) At equilibrium, the total pressure of the system is 0.390 atm. What is the partial pressure of each gas? A. 0.0303 atm H2; 0.360 atm H2O B. 0.0522 atm H2; 0.339 atm H2O C. 0.0696 atm H2; 0.320 atm H2O D. 0.320 atm H2; 0.0696 atm H2O E. 0.339 atm H2; 0.0522 atm H2O
27. At 25C, the decomposition of dinitrogen tetraoxide N2O4(g) 2NO2(g) has an equilibrium constant (K) of 0.144. At equilibrium, the total pressure of the system is 0.500 atm. What is the partial pressure of each gas? A. N2O4 = 0.206 atm; NO2 = 0.294 atm B. N2O4 = 0.212 atm; NO2 = 0.288 atm C. N2O4 = 0.288 atm; NO2 = 0.212 atm D. N2O4 = 0.294 atm; NO2 = 0.206 atm E. N2O4 = 0.437 atm; NO2 = 0.063 atm
28. At sufficiently high temperatures, ammonium iodide decomposes to ammonia and hydrogen iodide. NH4I(s) NH3(g) + HI(g) A mass of 5.00 g of NH4I is sealed in a 2.00-L flask and heated to 673 K. If 2.56 g NH4I(s) remain unreacted when the system has reach equilibrium, what is the equilibrium constant (Kp) for the reaction? (R = 0.0821 Latm/molK) A. 0.15 B. 0.22 C. 0.47 D. 0.59 E. 0.89
29. Hydrogen iodide can decompose into hydrogen and iodine gases. 2HI(g) H2(g) + I2(g) K for the reaction is 0.016. If 0.148 atm of HI(g) is sealed in a flask, what is the pressure of each gas when equilibrium is established? A. HI = 0.118 atm; H2 = 0.015 atm; I2 = 0.015 atm B. HI = 0.133 atm; H2 = 0.015 atm; I2 = 0.015 atm C. HI = 0.110 atm; H2 = 0.019 atm; I2 = 0.019 atm D. HI = 0.126 atm; H2 = 0.022 atm; I2 = 0.022 atm E. HI = 0.174 atm; H2 = 0.022 atm; I2 = 0.022 atm
30. Hydrogen iodide can decompose into hydrogen and iodine gases. 2HI(g) H2(g) + I2(g) Kp for the reaction is 0.016. If 0.350 atm of HI(g) is sealed in a flask, what is the total pressure of the system when equilibrium is established? A. 0.258 atm B. 0.279 atm C. 0.350 atm D. 0.385 atm E. 0.412 atm
31. Nitrosyl bromide decomposes according to the chemical equation below. 2NOBr(g) 2NO(g) + Br2(g) 1.00 atm of NOBr is sealed in a flask. At equilibrium, the partial pressure of NOBr is 0.82 atm. What is the equilibrium constant for the reaction? A. 3.6 ´ 10-3 B. 8.7 ´ 10-3 C. 2.8 ´ 10-2 D. 3.5 ´ 10-2 E. 4.3 ´ 10-3
32. Nitrosyl chloride decomposes according to the chemical equation below. 2NOCl(g) 2NO(g) + Cl2(g) A pressure of 0.320 atm of nitrosyl chloride is sealed in a flask and allowed to reach equilibrium. If 22.6% of the NOCl decomposes, what is the equilibrium constant for the reaction? A. 0.00153 B. 0.00308 C. 0.00611 D. 0.00730 E. 0.02471
33. At a given temperature, the equilibrium constant (Kp) for the decomposition of dinitrogen tetraoxide to nitrogen dioxide is 0.172. If 0.224 atm N2O4 is sealed in a flask, what partial pressure of NO2 will exist at equilibrium? N2O4(g) 2NO2(g) A. 0.0385 atm B. 0.158 atm C. 0.196 atm D. 0.257 atm E. 0.379 atm
34. The reaction of nitrogen gas and oxygen gas to form nitrogen monoxide, N2(g) + O2(g) 2NO(g) has an equilibrium constant of 1.0 ´ 10-30 at 298 K. What equilibrium partial pressure of NO(g) will form if 0.50 atm of N2 and 0.50 atm of O2 are sealed in a flask at 298 K? A. 1.0 ´ 10-60 atm B. 5.0 ´ 10-31 atm C. 1.0 ´ 10-30 atm D. 1.0 ´ 10-15 atm E. 5.0 ´ 10-16 atm
35. For the system CO(g) + H2O(g) CO2(g) + H2(g) K is 1.6 at 900 K. If 0.400 atm CO(g) and 0.400 atm H2O(g) are combined in a sealed flask, what is the equilibrium partial pressure of CO2(g)? A. 0.22 atm B. 0.31 atm C. 0.47 atm D. 0.51 atm E. 0.65 atm
36. Consider the following equilibrium: N2(g) + O2(g) 2NO(g) At a certain temperature the equilibrium constant for the reaction is 0.0255. What is the partial pressure of NO gas at equilibrium if the initial pressure of all the gases (both reactants and products) is 0.300 atm? A. 6.65 ´ 10-2 B. 0.183 C. 0.234 D. 0.252 E. 0.417
37. For the reaction 2A 3B Kc = 1.37. If the concentrations of A and B are equal, what is the value of that concentration? A. 0.685 M B. 0.822 M C. 1.17 M D. 1.37 M E. 1.88 M
38. Which of the following may change the ratio of products to reactants in an equilibrium mixture for a chemical reaction involving gaseous species? 1. 2. 3.
Increasing the temperature. Adding a catalyst. Adding gaseous reactants.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1 and 3 39. Assume that the following chemical reaction is at equilibrium. 2ICl(g)
I2(g) + Cl2(g)
DH = +26.9 kJ
At 25C, K = 2.0 ´ 105. If the temperature is increase to 45C, which statement applies?
A. K will decrease and the reaction will proceed in the backward direction. B. K will decrease and the reaction will proceed in the forward direction. C. K will remain unchanged and the reaction will proceed in the forward direction. D. K will increase and the reaction will proceed in the backward direction. E. K will increase and the reaction will proceed in the forward direction. 40. Assuming the reaction below is at equilibrium, which of the following changes will drive the reaction to the left? C(s) + O2(g)
1. 2. 3.
CO2(g)
Increasing the temperature. Adding O2(g). Removing C(s).
DH = -393.5 kJ/mol
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 2 and 3 41. The formation of ammonia from elemental nitrogen and hydrogen is an exothermic process. N2(g) + 3H2(g)
2NH3(g)
DH = -92.2 kJ/mol
Assuming the reaction is at equilibrium, which one of the following changes will drive the reaction to the right?
A. adding ammonia B. increasing the temperature C. increasing the pressure D. removing hydrogen E. adding a catalyst 42. Assume that the following endothermic chemical reaction is at equilibrium. C(s) + H2O(g) H2(g) + CO(g) Which of the following statements is/are CORRECT? 1. 2. 3.
Increasing the concentration of H2(g) will cause the reaction to proceed in the backward direction, increasing the equilibrium concentration of H2O(g). Decreasing the temperature will cause the reaction to proceed in the forward direction, increasing the equilibrium concentration of CO(g). Increasing the amount of C(s) will cause the reaction to proceed in the forward direction, increasing the equilibrium concentration of CO(g).
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 43. In which of the following equilibrium systems will an increase in the pressure have no effect on the concentrations of products and reactants? A. H2(g) + F2(g) 2 HF(g) B. N2(g) + 3 H2(g) 2 NH3(g) C. CaCO3(s) CaO(s) + CO2(g) D. 2 NOBr(g) 2 NO(g) + Br2(g) E. 2 H2O(g) + O2(g) 2 H2O2(g)
44. In which of the following equilibrium systems would an increase in volume (at constant temperature) cause the reaction to shift to the right? A. N2O4(g) 2NO2(g) B. N2(g) + 3H2(g) 2NH3(g) C. H2(g) + Cl2(g) 2HCl(g) D. Answers a and b are correct. E. Answers b and c are correct.
45. The equilibrium constant for a gas phase reaction is measured at two temperatures. At 100C, the equilibrium constant is 36. At 200C, the equilibrium constant is 147. Which of the following statements is correct for this equilibrium? A. The reaction must be first-order. B. The reaction is endothermic. C. A catalyst must be present. D. Each reactant molecule decomposes into two or more product molecules. E. One of the products must be a solid.
46. Which of the following statements is/are CORRECT for the following system at equilibrium? H2(g) + F2(g) 2 HF(g) 1. 2. 3.
Addition of a catalyst will increase the equilibrium constant, causing the reaction to proceed in the forward direction. The values of Kc and Kp are identical. Increasing the pressure of a system will cause the reaction to proceed in the forward direction.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 47. The equilibrium constant, Kp, for the reaction below is 0.24 at 1500C. SO3(g) + NO(g) SO2(g) + NO2(g) If 0.30 atm of sulfur trioxide, 0.15 atm of nitrogen monoxide, 0.55 atm of sulfur dioxide, and 0.030 atm of nitrogen dioxide are mixed, what changes in pressure will occur? A. The pressures of SO3 and NO decrease; the pressures of SO2 and NO2 increase. B. The pressures of SO3 and NO increase; the pressures of SO2 and NO2 decrease. C. The pressures of SO3 and SO2 decrease; the pressures of NO and NO2 increase. D. The pressures of SO3 and SO2 increase; the pressures of NO and NO2 decrease. E. Equal numbers of particles exist on both sides of the equation; no reaction will occur.
48. If the value of Q is less than Kp, then A. the system is in equilibrium. B. a catalyst is necessary to achieve equilibrium. C. the reaction will go left or right depending upon the reaction stoichiometry. D. the reaction will proceed to the right until equilibrium is established. E. the reaction will proceed to the left until equilibrium is established.
49. A 2.5 L flask is filled with 0.25 atm SO3, 0.20 atm SO2, and 0.40 atm O2, and allowed to reach equilibrium. Assume the temperature of the mixture is chosen so that Kp = 0.12. Predict the effect on the partial pressure of SO3 as equilibrium is achieved by using Q, the reaction quotient. 2 SO3(g) 2 SO2(g) + O2(g) A. The partial pressure of SO3 will decrease because Q > K. B. The partial pressure of SO3 will decrease because Q < K. C. The partial pressure of SO3 will increase because Q < K. D. The partial pressure of SO3 will increase because Q > K. E. The partial pressure of SO3 will remain the same because Q = K.
50. For the reaction C(s) + CO2(g) 2CO(g) Kc = 168. A mixture contains some C(s), [CO] = 0.50 M and [CO2] = 0.75 M. Therefore the system ____ at equilibrium, because ____. A. is not; the value of Q is 0.67 B. is not; the value of Q is 1.5 C. is; the value of Q is 0.67 D. is not; the value of Q is 0.33 E. is; the value of Q is 0.33
51. Consider the reaction A(g) 2B(g) where Kp = 4.1 at 25C. If 0.75 atm A(g) and 1.5 atm B(g) are initially present in a 1.0 L flask at 25C, what change in partial pressures (if any) will occur in time? A. The partial pressure of A will decrease and the partial pressure of B will decrease. B. The partial pressure of A will decrease and the partial pressure of B will increase. C. The partial pressure of A will increase and the partial pressure of B will decrease. D. The partial pressure of A will increase and the partial pressure of B will increase. E. The partial pressures of both A and B will remain unchanged.
52. "If a chemical system at equilibrium is disturbed by adding a gaseous species (reactant or product), the reaction will proceed in such a direction as to consume part of the added species" is a statement of A. the ideal gas law. B. Le Châtelier's principle. C. the de Broglie equation. D. the van't Hoff equation. E. the first law of thermodynamics.
53. Calcium carbonate decomposes to calcium oxide and carbon dioxide. CaCO3(s)
CaO(s) + CO2(g)
DH = +179 kJ
The equilibrium constant for this reaction is 9.7 ´ 10-24 at 298 K. What is the equilibrium constant at 575 K? (R = 8.31 J/molK)
A. 7.5 ´ 10-16 B. 1.3 ´ 10-8 C. 1.4 ´ 1038 D. 1.3 ´ 1015 E. 1.0 ´ 1023 54. The reaction of nitrogen with hydrogen to form ammonia is thermodynamically favorable. N2(g) + 3H2(g)
2NH3(g)
DH = -92.2 kJ
The equilibrium constant for this reaction is 6.0 ´ 105 at 298 K. At what temperature is the equilibrium constant equal to 1.0 ´ 103? (R = 8.31 J/molK)
A. 85 K B. 110 K C. 310 K D. 360 K E. 2800 K 55. The Haber process for the production of ammonia relies on high temperatures and pressures. Which of these, high temperatures or pressures, actually reduce the yield of the reaction at equilibrium? N2(g) + 3H2(g)
A. high pressure B. high temperature C. both D. neither E. can't be determined
2NH3(g)
DH = -92.2 kJ
56. The Haber process for the production of ammonia relies on a heterogeneous catalyst. How does the use of this catalyst effect the yield of the reaction at equilibrium? N2(g) + 3H2(g)
2NH3(g)
A. yield is increased B. yield is decreased C. depends on the catalyst used D. doesn't effect yield E. can't be determined
DH = -92.2 kJ
Chapter 13--Acids and Bases 1. According to the Brønsted-Lowry definition, a base A. is a weak electrolyte. B. increases the OH– concentration in an aqueous solution. C. is an electron-pair donor. D. increases the pH of a solution. E. is a proton acceptor.
2. Which of the following reactions are acid-base reactions according to the Brønsted-Lowry model? 1. 2. 3.
NH4+(aq) + H2O(l) 2HF(aq) + Ca2+(aq) H2PO4-(aq) + OH-(aq)
NH3(aq) + H3O+(aq) CaF2(s) + 2H+(aq) HPO42-(aq) + H 2O(l)
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 3. The species formed when a proton is removed from an acid is called the A. cation. B. conjugate base. C. conjugate acid. D. buffer. E. antacid.
4. A species that can either accept or donate a proton is called A. a Brønsted-Lowry compound. B. a Lewis base. C. an Arrhenius acid. D. amphiprotic. E. a conjugate pair.
5. What is the conjugate base of HClO2? A. H+ B. HC. ClO2+ D. ClO+ E. ClO2-
6. What is the conjugate acid of ammonia? A. H+ B. HC. NH4+ D. NH3 E. NH2-
7. What is the conjugate base of water? A. H+ B. O2C. OHD. H2O E. H3O+
8. What is the conjugate acid of potassium hydrogen phosphate, K2HPO4? A. H3PO4 B. H2PO4C. HPO42D. K+ E. OH-
9. All of the following species are amphiprotic EXCEPT A. H2O. B. HSO4-. C. HPO42-. D. SO42-. E. H2PO4-.
10. Which of the following equilibrium constant expressions represents the ionization of water? A.
B.
C. D. KW = -log[H+] E. KW = -log[OH-]
11. Which of the following chemical reactions represents the ionization of water? A. 2H2O(l) 2H2(g) + O2(g) B. H2O(l) H2O-(aq) C. H2O(l) H2O+(aq) + eD. H2O(l) H+(aq) + OH-(aq) E. H+(aq) + OH-(aq) H2O(l)
12. If the H+ concentration in a carbonated beverage is 4.7 ´ 10-5 M, what is the OH- concentration? (Kw = 1.0 ´ 10-14) A. 4.7 ´ 10-19 M B. 2.1 ´ 10-10 M C. 5.3 ´ 10-5 M D. 2.1 ´ 104 M E. 4.7 ´ 109 M
13. If the OH- concentration in a bottle of an ammonia based cleaner is 3.8 ´ 10-2 M, what is the H+ concentration? (Kw = 1.0 ´ 10-14) A. 3.8 ´ 10-16 M B. 2.6 ´ 10-13 M C. 6.8 ´ 10-2 M D. 2.6 ´ 101 M E. 3.8 ´ 1012 M
14. If the H+ concentration is less than 1.0 ´ 10-7 M, the solution is A. basic. B. acidic. C. neutral. D. amphiprotic. E. in equilibrium.
15. What is the correct expression for the calculation of pH? A. pH = log[1.0 ´ 10-14] B. pH = -log([H+][OH-]) C. pH = log[OH-] D. pH = -log[OH-] E. pH = -log[H+]
16. If the pH of a solution is greater than 7, the solution is A. acidic. B. basic. C. amphiprotic. D. neutral. E. in equilibrium.
17. What is the H3O+ concentration of an aqueous solution with a pH of 8.77? A. 1.7 ´ 10–9 M B. 5.9 ´ 10–6 M C. 1.6 ´ 10–4 M D. 5.23 M E. 5.9 ´ 108 M
18. The pH of a human blood sample is 7.30. What is concentration of OH- in blood? A. 5.01 ´ 10-8 M B. 2.0 ´ 10-7 M C. 7.3 ´ 10-7 M D. 5.01 ´ 10-5 M E. 2.0 ´ 107 M
19. The hydronium ion concentration in a sample of lemon juice is 3.2 ´ 10-3 M. What is the pH of the lemon juice? A. 1.06 B. 2.01 C. 2.49 D. 5.21 E. 11.72
20. Seawater has a hydroxide ion concentration of 2.0 ´ 10-6 M. What is the pH of seawater? A. -8.30 B. 5.70 C. 6.99 D. 7.53 E. 8.30
21. What is the pH of 6.5 ´ 10–5 M KOH(aq) at 25C? A. –4.19 B. 1.54 C. 4.19 D. 9.81 E. 12.46
22. All of the following species are strong acids EXCEPT A. HClO4. B. HBr. C. H2SO4. D. HF. E. HI.
23. All of the following species are strong bases EXCEPT A. NaOH. B. KOH. C. Mg(OH)2. D. Sr(OH)2. E. RbOH.
24. Which of the following species is a weak base? A. KOH B. NH4+ C. HF D. H3PO4 E. CH3CO2-
25. What is the pH of a solution prepared by diluting 0.40 mol HNO3(aq) to a volume of 225 mL? A. -0.25 B. 0.25 C. 0.40 D. 2.75 E. 7.50
26. An aqueous solution with a pH of 2.00 is diluted from 1.0 L to 3.0 L. What is the pH of the diluted solution? A. 0.67 B. 2.00 C. 2.48 D. 4.33 E. 6.00
27. Which of the following solutions will have a pH of 3.0? A. 1 ´ 10-3 M CH3CO2H B. 1 ´ 10-3 M NH3 C. 1 ´ 10-3 M NH4+ D. 1 ´ 10-3 M HI E. Answers b and c are correct.
28. Which of the following solutions will have a pH of 11.0? A. 11 M Sr(OH)2 B. 1 ´ 10-11 M NH3 C. 1 ´ 10-11 M HCl D. 1 ´ 10-3 M NH4+ E. 1 ´ 10-3 M NaOH
29. Which of the following chemical reactions corresponds to the generic form of the base dissociation constant (Kb)? A. HB(aq) + OH-(aq) B-(aq) + H2O(l) B. 2H2O(l) H3O+(aq) + OH-(aq) C. HB(aq) + H2O(l) B-(aq) + H3O+(aq) D. B (aq) + H2O(l) HB(aq) + OH-(aq) E. B-(aq) + H3O+(aq) HB(aq) + H2O(l)
30. The acid equilibrium constant (Ka) for ammonium chloride refers to which of the following chemical reactions? A. NH4+(aq) + H3O+(aq) NH52+(aq) + H2O(l) + B. NH4 (aq) + H2O(l) NH3(aq) + H3O+(aq) C. NH4+(aq) + OH-(aq) NH3(aq) + H2O(l) D. NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) E. NH3(aq) + H3O+(aq) NH4+(aq) + H2O(l)
31. Which of the following chemical equations corresponds to the base ionization constant, Kb, for hydrogen sulfite ion (HSO3–)? A. HSO3–(aq) + H2SO3(aq) 2 H2SO3(aq) B. HSO3–(aq) + H2O(l) SO32–(aq) + H3O+(aq) C. HSO3–(aq) + OH–(aq) SO32–(aq) + H2O(l) D. HSO3–(aq) + H3O+(aq) H2SO3(aq) + H2O(l) E. HSO3–(aq) + H2O(l) H2SO3(aq) + OH–(aq)
32. Which of the following mathematical equations corresponds to the acid dissociation constant (Ka)?
A.
B.
C.
D.
E.
33. The anions derived from strong acids are A. spectator ions. B. strong bases. C. amphiprotic. D. weak acids. E. strong acids.
34. The conjugate acid of a weak base is A. amphiprotic. B. a weak base. C. a weak acid. D. a strong acid. E. a strong base.
35. At 50C, the water ionization constant, Kw, is 5.48 ´ 10–14. What is the H3O+ concentration in neutral water at this temperature? A. 3.00 ´ 10–27 M B. 2.74 ´ 10–14 M C. 5.48 ´ 10–14 M D. 2.34 ´ 10–7 M E. 1.01 ´ 10–7 M
36. What is the conjugate base of [Fe(H2O)6]3+(aq)? A. H3O+ B. [Fe(H2O)6]2+ C. [Fe(H2O)5H3O]4+ D. [Fe(H2O)5OH]2+ E. [Fe(H2O)5]3+
37. A solution of 0.10 M aluminum nitrate has a pH close to 3.0. Which chemical equation explains the acidic pH of the solution? A. Al3+(aq) + 3OH-(aq) Al(OH)3(s) + 3H3O+(aq) 3+ B. Al(H2O)6 (aq) + H2O(l) Al(H2O)5(OH)2+(aq) + H3O+(aq) C. Al(NO3)3(s) + 3H2O(l) Al(OH)3(s) + 3HNO3(s) D. HNO3(aq) + H2O(l) NO3-(aq) + H3O+(aq) E. NO3-(aq) + H2O(l) H2O + HNO3(aq)
38. The pH of aqueous 0.50 M hypobromous acid, HBrO, is 4.45. What is the Ka of this acid? A. 2.5 ´ 10–9 B. 5.0 ´ 10–9 C. 3.4 ´ 10–7 D. 3.5 ´ 10–5 E. 7.1 ´ 10–5
39. The pH of 0.350 M benzoic acid, HC7H5O2, is 2.32. What is the pKa for this acid? A. 0.350 B. 1.38 C. 1.79 D. 2.32 E. 4.18
40. The pH of 0.400 M sodium nitrite, NaNO2, is 8.42. What is the Kb for this base? A. 1.7 ´ 10-11 B. 3.6 ´ 10-10 C. 3.8 ´ 10-9 D. 9.5 ´ 10-9 E. 8.8 ´ 10-7
41. The pH of aqueous 0.10 M pyridine (C5H5N) ion is 9.09. What is the Kb of this base? A. 8.0 ´ 10–10 B. 1.5 ´ 10–9 C. 9.0 ´ 10–6 D. 1.6 ´ 10–5 E. 1.2 ´ 10–5
42. A solution of 0.25 mol HF diluted to 1.0 L has a hydronium ion concentration of 1.3 ´ 10-2 M. What is the percent ionization of HF? A. 0.80 % B. 1.3% C. 4.2% D. 5.2% E. 31%
43. What is the pH of 0.26 M ammonium ion? NH4+(aq) + H2O(l)
NH3(aq) + H3O+(aq)
Ka = 5.6 ´ 10-10
A. 3.87 B. 4.33 C. 4.75 D. 4.92 E. 9.25 44. Ammonia is a weak base (Kb = 1.8 ´ 10-5). What is the pH of 1.2 M ammonia? A. 2.33 B. 4.74 C. 9.26 D. 10.14 E. 11.67
45. Lactic acid, HC3H5O3, is found in sour milk. What is the pH of 0.30 M lactic acid? (Ka = 1.4 ´ 10-4) A. 0.52 B. 1.07 C. 2.19 D. 4.00 E. 5.12
46. Hydrogen sulfate ion, HSO4-, has an acid dissociation constant of 1.0 ´ 10-2. What is the pH of 0.45 M hydrogen sulfate? A. 0.35 B. 1.21 C. 2.00 D. 2.19 E. 3.03
47. What is the pH of 0.25 M aqueous acetate ion? (Kb of CH3CO2– = 5.6 ´ 10–10) A. 4.32 B. 4.93 C. 9.07 D. 9.68 E. 13.40
48. Which of the following chemical equations corresponds to Ka3 for phosphoric acid, H3PO4? A. H3PO4(aq) + H2O(l) H2PO4-(aq) + H3O+(aq) B. H2PO4 (aq) + H2O(l) HPO42-(aq) + H3O+(aq) C. HPO42-(aq) + H2O(l) PO43-(aq) + H3O+(aq) D. H3PO4(aq) + 2H2O(l) HPO42-(aq) + 2H3O+(aq) E. H2PO4 (aq) + 2H2O(l) PO43-(aq) + H3O+(aq)
49. Which of the following chemical equations corresponds to Kb1 for SO32–? A. HSO3–(aq) + H2O(l) H2SO3(aq) + OH–(aq) 2– + B. SO3 (aq) + H3O (aq) HSO3–(aq) + H2O(l) C. H2SO3(aq) + OH–(aq) HSO3–(aq) + H2O(l) – – D. HSO3 (aq) + OH (aq) SO32–(aq) + H2O(l) E. SO32–(aq) + H2O(l) HSO3–(aq) + OH–(aq)
50. All of the following statements concerning diprotic acids and bases are incorrect EXCEPT A. Ka1 is always large. B. Ka1 is larger than Ka2. C. Ka1 is equal to Kb2. D. Ka1 + Kb2 equals Kw. E. Ka1 + Ka2 equals Kw.
51. What is the pH of 1.0 M sulfurous acid? (Ka1 = 1.7 ´ 10-2, Ka2 = 6.0 ´ 10-8) A. -0.88 B. 0.00 C. 0.12 D. 0.13 E. 0.91
52. The Ka for hydrofluoric acid is 6.9 ´ 10-4. What is Kb for fluoride ion? A. 1.0 ´ 10-14 B. 1.4 ´ 10-11 C. 7.4 ´ 10-9 D. 1.0 ´ 10-7 E. 6.9 ´ 1010
53. Given the following equilibrium constants, Ka (HSO4–) = 1.2 ´ 10–2 Kb (CH3CO2–) = 5.6 ´ 10–10 Kw = 1.00 ´ 10–14 determine the equilibrium constant for the reaction below at 25C. HSO4–(aq) + CH3CO2–(aq) SO42–(aq) + CH3CO2H(aq) –12 A. 6.7 ´ 10 B. 2.1 ´ 10–7 C. 1.5 ´ 10–3 D. 6.7 ´ 102 E. 2.1 ´ 107
54. Determine the equilibrium constant for the reaction HF(aq) + NH3(aq) F-(aq) + NH4+(aq) given the equilibrium constants for the following reactions. HF(aq) + H2O(l) F-(aq) + H3O +(aq) NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) 2H2O OH-(aq) + H3O+(aq)
Ka = 6.9 ´ 10-4 Kb = 1.8 ´ 10-5 Kw = 1.0 ´ 10-14
A. 1.2 ´ 10-8 B. 1.2 ´ 106 C. 8.1 ´ 107 D. 1.0 ´ 1014 E. 3.8 ´ 1015 55. A salt solution can be acidic, basic, or neutral. When dissolved in water, which of the following salts will make the solution acidic: NaCl, Al2(SO4)3, NaNO3, Na2CO3, KF, and NH4Br? A. Al2(SO4)3 and NH4Br B. Al2(SO4)3, Na2CO3 and NH4Br C. Al2(SO4)3 and NaNO3 D. NaCl, NaNO3 and NH4Br E. Na2CO3, KF, and NH4Br
56. A salt solution can be acidic, basic, or neutral. When dissolved in water, which of the following salts will make the solution basic: KBr, MgCO3, NaF, Na2SO4, and CrCl3? A. MgCO3 and CrCl3 B. KBr and Na2SO4 C. KBr, MgCO3, and CrCl3 D. MgCO3, NaF, and Na2SO4 E. NaF, Na2SO4, and CrCl3
57. A salt solution can be acidic, basic, or neutral. When dissolved in water, which of the following salts will not affect the pH: KCl, FeCl3, NaNO3, CaCO3, LiF, and NH4Br? A. KCl and FeCl3 B. KCl, FeCl3 and NaNO3 C. CaCO3 and LiF D. KCl and NaNO3 E. CaCO3, LiF, and NH4Br
58. All carboxylic acids contain the organic group(s) ______ . A. -CH3 B. -NH2 C. -OH D. -COOH E. both -NH2 and -COOH
59. A Lewis acid is defined as a(n) A. electron pair acceptor. B. proton acceptor. C. electron pair donor. D. spectator ion. E. proton acceptor.
60. Identify from the following list of molecules and ions those which behave as Lewis acids: NH3, BCl3, Fe3+. A. NH3 only B. NH3 and BCl3 C. Fe3+ only D. BCl3, and Fe3+ E. NH3, BCl3, and Fe3+
Chapter 13--Acids and Bases Key
1. According to the Brønsted-Lowry definition, a base A. is a weak electrolyte. B. increases the OH– concentration in an aqueous solution. C. is an electron-pair donor. D. increases the pH of a solution. E. is a proton acceptor.
2. Which of the following reactions are acid-base reactions according to the Brønsted-Lowry model? 1. 2. 3.
NH4+(aq) + H2O(l) 2HF(aq) + Ca2+(aq) H2PO4-(aq) + OH-(aq)
NH3(aq) + H3O+(aq) CaF2(s) + 2H+(aq) HPO42-(aq) + H 2O(l)
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 3. The species formed when a proton is removed from an acid is called the A. cation. B. conjugate base. C. conjugate acid. D. buffer. E. antacid.
4. A species that can either accept or donate a proton is called A. a Brønsted-Lowry compound. B. a Lewis base. C. an Arrhenius acid. D. amphiprotic. E. a conjugate pair.
5. What is the conjugate base of HClO2? A. H+ B. HC. ClO2+ D. ClO+ E. ClO2-
6. What is the conjugate acid of ammonia? A. H+ B. HC. NH4+ D. NH3 E. NH2-
7. What is the conjugate base of water? A. H+ B. O2C. OHD. H2O E. H3O+
8. What is the conjugate acid of potassium hydrogen phosphate, K2HPO4? A. H3PO4 B. H2PO4C. HPO42D. K+ E. OH-
9. All of the following species are amphiprotic EXCEPT A. H2O. B. HSO4-. C. HPO42-. D. SO42-. E. H2PO4-.
10. Which of the following equilibrium constant expressions represents the ionization of water? A.
B.
C. D. KW = -log[H+] E. KW = -log[OH-]
11. Which of the following chemical reactions represents the ionization of water? A. 2H2O(l) 2H2(g) + O2(g) B. H2O(l) H2O-(aq) C. H2O(l) H2O+(aq) + eD. H2O(l) H+(aq) + OH-(aq) E. H+(aq) + OH-(aq) H2O(l)
12. If the H+ concentration in a carbonated beverage is 4.7 ´ 10-5 M, what is the OH- concentration? (Kw = 1.0 ´ 10-14) A. 4.7 ´ 10-19 M B. 2.1 ´ 10-10 M C. 5.3 ´ 10-5 M D. 2.1 ´ 104 M E. 4.7 ´ 109 M
13. If the OH- concentration in a bottle of an ammonia based cleaner is 3.8 ´ 10-2 M, what is the H+ concentration? (Kw = 1.0 ´ 10-14) A. 3.8 ´ 10-16 M B. 2.6 ´ 10-13 M C. 6.8 ´ 10-2 M D. 2.6 ´ 101 M E. 3.8 ´ 1012 M
14. If the H+ concentration is less than 1.0 ´ 10-7 M, the solution is A. basic. B. acidic. C. neutral. D. amphiprotic. E. in equilibrium.
15. What is the correct expression for the calculation of pH? A. pH = log[1.0 ´ 10-14] B. pH = -log([H+][OH-]) C. pH = log[OH-] D. pH = -log[OH-] E. pH = -log[H+]
16. If the pH of a solution is greater than 7, the solution is A. acidic. B. basic. C. amphiprotic. D. neutral. E. in equilibrium.
17. What is the H3O+ concentration of an aqueous solution with a pH of 8.77? A. 1.7 ´ 10–9 M B. 5.9 ´ 10–6 M C. 1.6 ´ 10–4 M D. 5.23 M E. 5.9 ´ 108 M
18. The pH of a human blood sample is 7.30. What is concentration of OH- in blood? A. 5.01 ´ 10-8 M B. 2.0 ´ 10-7 M C. 7.3 ´ 10-7 M D. 5.01 ´ 10-5 M E. 2.0 ´ 107 M
19. The hydronium ion concentration in a sample of lemon juice is 3.2 ´ 10-3 M. What is the pH of the lemon juice? A. 1.06 B. 2.01 C. 2.49 D. 5.21 E. 11.72
20. Seawater has a hydroxide ion concentration of 2.0 ´ 10-6 M. What is the pH of seawater? A. -8.30 B. 5.70 C. 6.99 D. 7.53 E. 8.30
21. What is the pH of 6.5 ´ 10–5 M KOH(aq) at 25C? A. –4.19 B. 1.54 C. 4.19 D. 9.81 E. 12.46
22. All of the following species are strong acids EXCEPT A. HClO4. B. HBr. C. H2SO4. D. HF. E. HI.
23. All of the following species are strong bases EXCEPT A. NaOH. B. KOH. C. Mg(OH)2. D. Sr(OH)2. E. RbOH.
24. Which of the following species is a weak base? A. KOH B. NH4+ C. HF D. H3PO4 E. CH3CO2-
25. What is the pH of a solution prepared by diluting 0.40 mol HNO3(aq) to a volume of 225 mL? A. -0.25 B. 0.25 C. 0.40 D. 2.75 E. 7.50
26. An aqueous solution with a pH of 2.00 is diluted from 1.0 L to 3.0 L. What is the pH of the diluted solution? A. 0.67 B. 2.00 C. 2.48 D. 4.33 E. 6.00
27. Which of the following solutions will have a pH of 3.0? A. 1 ´ 10-3 M CH3CO2H B. 1 ´ 10-3 M NH3 C. 1 ´ 10-3 M NH4+ D. 1 ´ 10-3 M HI E. Answers b and c are correct.
28. Which of the following solutions will have a pH of 11.0? A. 11 M Sr(OH)2 B. 1 ´ 10-11 M NH3 C. 1 ´ 10-11 M HCl D. 1 ´ 10-3 M NH4+ E. 1 ´ 10-3 M NaOH
29. Which of the following chemical reactions corresponds to the generic form of the base dissociation constant (Kb)? A. HB(aq) + OH-(aq) B-(aq) + H2O(l) B. 2H2O(l) H3O+(aq) + OH-(aq) C. HB(aq) + H2O(l) B-(aq) + H3O+(aq) D. B (aq) + H2O(l) HB(aq) + OH-(aq) E. B-(aq) + H3O+(aq) HB(aq) + H2O(l)
30. The acid equilibrium constant (Ka) for ammonium chloride refers to which of the following chemical reactions? A. NH4+(aq) + H3O+(aq) NH52+(aq) + H2O(l) + B. NH4 (aq) + H2O(l) NH3(aq) + H3O+(aq) C. NH4+(aq) + OH-(aq) NH3(aq) + H2O(l) D. NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) E. NH3(aq) + H3O+(aq) NH4+(aq) + H2O(l)
31. Which of the following chemical equations corresponds to the base ionization constant, Kb, for hydrogen sulfite ion (HSO3–)? A. HSO3–(aq) + H2SO3(aq) 2 H2SO3(aq) B. HSO3–(aq) + H2O(l) SO32–(aq) + H3O+(aq) C. HSO3–(aq) + OH–(aq) SO32–(aq) + H2O(l) D. HSO3–(aq) + H3O+(aq) H2SO3(aq) + H2O(l) E. HSO3–(aq) + H2O(l) H2SO3(aq) + OH–(aq)
32. Which of the following mathematical equations corresponds to the acid dissociation constant (Ka)?
A.
B.
C.
D.
E.
33. The anions derived from strong acids are A. spectator ions. B. strong bases. C. amphiprotic. D. weak acids. E. strong acids.
34. The conjugate acid of a weak base is A. amphiprotic. B. a weak base. C. a weak acid. D. a strong acid. E. a strong base.
35. At 50C, the water ionization constant, Kw, is 5.48 ´ 10–14. What is the H3O+ concentration in neutral water at this temperature? A. 3.00 ´ 10–27 M B. 2.74 ´ 10–14 M C. 5.48 ´ 10–14 M D. 2.34 ´ 10–7 M E. 1.01 ´ 10–7 M
36. What is the conjugate base of [Fe(H2O)6]3+(aq)? A. H3O+ B. [Fe(H2O)6]2+ C. [Fe(H2O)5H3O]4+ D. [Fe(H2O)5OH]2+ E. [Fe(H2O)5]3+
37. A solution of 0.10 M aluminum nitrate has a pH close to 3.0. Which chemical equation explains the acidic pH of the solution? A. Al3+(aq) + 3OH-(aq) Al(OH)3(s) + 3H3O+(aq) 3+ B. Al(H2O)6 (aq) + H2O(l) Al(H2O)5(OH)2+(aq) + H3O+(aq) C. Al(NO3)3(s) + 3H2O(l) Al(OH)3(s) + 3HNO3(s) D. HNO3(aq) + H2O(l) NO3-(aq) + H3O+(aq) E. NO3-(aq) + H2O(l) H2O + HNO3(aq)
38. The pH of aqueous 0.50 M hypobromous acid, HBrO, is 4.45. What is the Ka of this acid? A. 2.5 ´ 10–9 B. 5.0 ´ 10–9 C. 3.4 ´ 10–7 D. 3.5 ´ 10–5 E. 7.1 ´ 10–5
39. The pH of 0.350 M benzoic acid, HC7H5O2, is 2.32. What is the pKa for this acid? A. 0.350 B. 1.38 C. 1.79 D. 2.32 E. 4.18
40. The pH of 0.400 M sodium nitrite, NaNO2, is 8.42. What is the Kb for this base? A. 1.7 ´ 10-11 B. 3.6 ´ 10-10 C. 3.8 ´ 10-9 D. 9.5 ´ 10-9 E. 8.8 ´ 10-7
41. The pH of aqueous 0.10 M pyridine (C5H5N) ion is 9.09. What is the Kb of this base? A. 8.0 ´ 10–10 B. 1.5 ´ 10–9 C. 9.0 ´ 10–6 D. 1.6 ´ 10–5 E. 1.2 ´ 10–5
42. A solution of 0.25 mol HF diluted to 1.0 L has a hydronium ion concentration of 1.3 ´ 10-2 M. What is the percent ionization of HF? A. 0.80 % B. 1.3% C. 4.2% D. 5.2% E. 31%
43. What is the pH of 0.26 M ammonium ion? NH4+(aq) + H2O(l)
NH3(aq) + H3O+(aq)
Ka = 5.6 ´ 10-10
A. 3.87 B. 4.33 C. 4.75 D. 4.92 E. 9.25 44. Ammonia is a weak base (Kb = 1.8 ´ 10-5). What is the pH of 1.2 M ammonia? A. 2.33 B. 4.74 C. 9.26 D. 10.14 E. 11.67
45. Lactic acid, HC3H5O3, is found in sour milk. What is the pH of 0.30 M lactic acid? (Ka = 1.4 ´ 10-4) A. 0.52 B. 1.07 C. 2.19 D. 4.00 E. 5.12
46. Hydrogen sulfate ion, HSO4-, has an acid dissociation constant of 1.0 ´ 10-2. What is the pH of 0.45 M hydrogen sulfate? A. 0.35 B. 1.21 C. 2.00 D. 2.19 E. 3.03
47. What is the pH of 0.25 M aqueous acetate ion? (Kb of CH3CO2– = 5.6 ´ 10–10) A. 4.32 B. 4.93 C. 9.07 D. 9.68 E. 13.40
48. Which of the following chemical equations corresponds to Ka3 for phosphoric acid, H3PO4? A. H3PO4(aq) + H2O(l) H2PO4-(aq) + H3O+(aq) B. H2PO4 (aq) + H2O(l) HPO42-(aq) + H3O+(aq) C. HPO42-(aq) + H2O(l) PO43-(aq) + H3O+(aq) D. H3PO4(aq) + 2H2O(l) HPO42-(aq) + 2H3O+(aq) E. H2PO4 (aq) + 2H2O(l) PO43-(aq) + H3O+(aq)
49. Which of the following chemical equations corresponds to Kb1 for SO32–? A. HSO3–(aq) + H2O(l) H2SO3(aq) + OH–(aq) 2– + B. SO3 (aq) + H3O (aq) HSO3–(aq) + H2O(l) C. H2SO3(aq) + OH–(aq) HSO3–(aq) + H2O(l) – – D. HSO3 (aq) + OH (aq) SO32–(aq) + H2O(l) E. SO32–(aq) + H2O(l) HSO3–(aq) + OH–(aq)
50. All of the following statements concerning diprotic acids and bases are incorrect EXCEPT A. Ka1 is always large. B. Ka1 is larger than Ka2. C. Ka1 is equal to Kb2. D. Ka1 + Kb2 equals Kw. E. Ka1 + Ka2 equals Kw.
51. What is the pH of 1.0 M sulfurous acid? (Ka1 = 1.7 ´ 10-2, Ka2 = 6.0 ´ 10-8) A. -0.88 B. 0.00 C. 0.12 D. 0.13 E. 0.91
52. The Ka for hydrofluoric acid is 6.9 ´ 10-4. What is Kb for fluoride ion? A. 1.0 ´ 10-14 B. 1.4 ´ 10-11 C. 7.4 ´ 10-9 D. 1.0 ´ 10-7 E. 6.9 ´ 1010
53. Given the following equilibrium constants, Ka (HSO4–) = 1.2 ´ 10–2 Kb (CH3CO2–) = 5.6 ´ 10–10 Kw = 1.00 ´ 10–14 determine the equilibrium constant for the reaction below at 25C. HSO4–(aq) + CH3CO2–(aq) SO42–(aq) + CH3CO2H(aq) –12 A. 6.7 ´ 10 B. 2.1 ´ 10–7 C. 1.5 ´ 10–3 D. 6.7 ´ 102 E. 2.1 ´ 107
54. Determine the equilibrium constant for the reaction HF(aq) + NH3(aq) F-(aq) + NH4+(aq) given the equilibrium constants for the following reactions. HF(aq) + H2O(l) F-(aq) + H3O +(aq) NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) 2H2O OH-(aq) + H3O+(aq)
Ka = 6.9 ´ 10-4 Kb = 1.8 ´ 10-5 Kw = 1.0 ´ 10-14
A. 1.2 ´ 10-8 B. 1.2 ´ 106 C. 8.1 ´ 107 D. 1.0 ´ 1014 E. 3.8 ´ 1015 55. A salt solution can be acidic, basic, or neutral. When dissolved in water, which of the following salts will make the solution acidic: NaCl, Al2(SO4)3, NaNO3, Na2CO3, KF, and NH4Br? A. Al2(SO4)3 and NH4Br B. Al2(SO4)3, Na2CO3 and NH4Br C. Al2(SO4)3 and NaNO3 D. NaCl, NaNO3 and NH4Br E. Na2CO3, KF, and NH4Br
56. A salt solution can be acidic, basic, or neutral. When dissolved in water, which of the following salts will make the solution basic: KBr, MgCO3, NaF, Na2SO4, and CrCl3? A. MgCO3 and CrCl3 B. KBr and Na2SO4 C. KBr, MgCO3, and CrCl3 D. MgCO3, NaF, and Na2SO4 E. NaF, Na2SO4, and CrCl3
57. A salt solution can be acidic, basic, or neutral. When dissolved in water, which of the following salts will not affect the pH: KCl, FeCl3, NaNO3, CaCO3, LiF, and NH4Br? A. KCl and FeCl3 B. KCl, FeCl3 and NaNO3 C. CaCO3 and LiF D. KCl and NaNO3 E. CaCO3, LiF, and NH4Br
58. All carboxylic acids contain the organic group(s) ______ . A. -CH3 B. -NH2 C. -OH D. -COOH E. both -NH2 and -COOH
59. A Lewis acid is defined as a(n) A. electron pair acceptor. B. proton acceptor. C. electron pair donor. D. spectator ion. E. proton acceptor.
60. Identify from the following list of molecules and ions those which behave as Lewis acids: NH3, BCl3, Fe3+. A. NH3 only B. NH3 and BCl3 C. Fe3+ only D. BCl3, and Fe3+ E. NH3, BCl3, and Fe3+
Chapter 14--Equilibria in Acid-Base Solutions 1. An acid-base equilibrium system is created by dissolving 0.50 mol CH3CO2H in water to a volume of 1.0 L. What is the effect of adding 0.50 mol CH3CO2–(aq) to this solution? 1. 2. 3.
The pH of the solution will equal 7.00 because equal concentrations of a weak acid and its conjugate base are present. Some CH3CO2H(aq) will ionize, increasing the concentration of CH3CO2–(aq) and increasing the pH. Some CH3CO2–(aq) will react with H3O+(aq), increasing the concentration of CH3CO2H(aq) and reestablishing the solution equilibrium.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 2. What is the effect of adding NaOH(aq) to an aqueous solution of ammonia? 1. 2. 3.
The pH of the solution will increase. The concentration of NH4+(aq) will decrease. The concentration of NH3(aq) will increase.
A. 1 only B. 2 only C. 3 only D. 2 and 3 E. 1, 2, and 3 3. All of the following statements concerning acid-base buffers are true EXCEPT A. buffers are resistant to pH changes upon addition of small quantities of strong acids or bases. B. buffers are used as colored indicators in acid-base titrations. C. the pH of a buffer is close to the pKa of the weak acid from which it is made. D. buffers contain appreciable quantities of a weak acid and its conjugate base. E. buffers are resistant to changes in pH when diluted with water.
4. Which of the following pairs will form a buffer when mixed together in an aqueous solution? A. KCl and KH2PO4 B. HCl and KOH C. Ca(OH)2 and NaOH D. HF and NaF E. None of the above will form a buffer.
5. When mixed together, all of the following pairs can form buffers EXCEPT A. H3PO4 and NaH2PO4. B. NaH2PO4 and Na2HPO4. C. CH3CO2H and NaOH. D. HCl and NaCH3CO2. E. NaI and NaOH.
6. Which of the following equations is the Henderson-Hasselbalch equation?
A.
B.
C.
D.
E.
7. What is the pH of a solution that results from diluting 0.30 mol acetic acid (CH3CO2H) and 0.20 mol sodium acetate (NaCH3CO2) with water to a volume of 1.0 L? (Ka of CH3CO2H = 1.8 ´ 10–5) A. 4.35 B. 4.57 C. 4.74 D. 4.92 E. 5.14
8. What is the pH of an aqueous solution composed of 0.64 M NH4+ and 0.20 M NH3? (Ka of NH4+ = 5.6 ´ 10–10) A. 4.80 B. 8.75 C. 9.20 D. 9.25 E. 9.76
9. The Ka of hypobromous acid, HOBr, is 2.6 ´ 10-9. Calculate the pH of a solution which is composed of 0.40 M HOBr and 0.40 M NaOBr. A. 0.40 B. 0.80 C. 4.49 D. 8.59 E. 9.12
10. The Ka of acetic acid is 1.8 ´ 10-5. Calculate the pH of a solution that is composed of 1.0 M CH3CO2H and 0.50 M NaCH3CO2. A. 2.24 B. 4.44 C. 4.74 D. 5.05 E. 6.99
11. A buffer may be prepared by mixing a weak acid with a roughly equivalent amount of strong base. Which of the acids below is best for the preparation of a buffer with a pH of 4.00? A. sulfurous acid , H2SO3; Ka = 1.7 ´ 10-2 B. hydrofluoric acid, HF; Ka = 6.9 ´ 10-4 C. benzoic acid, HC7H5O2; Ka = 6.6 ´ 10-5 D. dihydrogen phosphate ion, H2PO4-; Ka = 6. 2 ´ 10-8 E. ammonium ion, NH4+; Ka = 5.6 ´ 10-10
12. A buffer may be prepared by mixing a weak acid with a roughly equivalent amount of strong base. Which of the acids below is best for the preparation of a buffer with a pH of 9.00? A. chlorous acid, HClO2; Ka = 2.8 ´ 10-8 B. formic acid, HCO2H; Ka = 1.9 ´ 10-4 C. benzoic acid, HC7H5O2; Ka = 6.6 ´ 10-5 D. dihydrogen phosphate ion, H2PO4-; Ka = 6.2 ´ 10-8 E. ammonium ion, NH4+; Ka = 5.6 ´ 10-10
13. What is the pH of the buffer that results when 12.0 g of NaH2PO4 and 8.00 g of Na2HPO4 are diluted with water to a volume of 0.50 L? (Ka of H2PO4– = 6.2 ´ 10–8, the molar masses of NaH2PO4 and Na2HPO4 are 120.0 g/mol and 142.0 mol, respectively) A. 4.10 B. 6.96 C. 7.21 D. 7.46 E. 9.90
14. An acetic acid-sodium acetate buffer is prepared by mixing 6.00 g NaCH3CO2 with 4.00 mL of CH3CO2H (density = 1.042 g/mL) and diluting to 1.00 L. What is the pH of the resulting buffer? The Ka of acetic acid is 1.8 ´ 10-5. A. 4.10 B. 4.77 C. 4.99 D. 5.19 E. 5.78
15. What is the pH of the buffer that results when 32 g sodium acetate (NaCH3CO2) is mixed with 500.0 mL of 1.0 M acetic acid (CH3CO2H) and diluted with water to 1.0 L? (Ka of CH3CO2H = 1.8 ´ 10–5) A. 2.52 B. 4.23 C. 4.44 D. 4.64 E. 4.74
16. A buffer is prepared by combining 25 mL of 0.50 M NaF(aq) with 25 mL of 0.25 M HCl. What is the pH of the buffer? (Ka (HF) = 6.9 ´ 10-4) A. 3.16 B. 3.50 C. 4.12 D. 4.60 E. 7.11
17. A buffer is prepared by combining 25 mL of 0.50 M NH3(aq) with 25 mL of 0.20 M HCl. What is the pH of the buffer? (Ka (NH4+) = 5.6 ´ 10-10) A. 7.76 B. 8.00 C. 8.86 D. 9.43 E. 9.65
18. A buffer is prepared by combining 10.0 g of NaH2PO4 with 150 mL of 0.20 M NaOH and diluting to 2.0 L. What is the pH of the buffer? (Ka (H2PO4-) = 6.2 ´ 10-8) A. 5.12 B. 6.76 C. 6.96 D. 7.20 E. 7.45
19. What is the pH of a buffer that results when 0.50 mole of H3PO4 is mixed with 0.75 mole of NaOH and diluted with water to 1.00 L? (The acid dissociation constants of phosphoric acid are Ka1 = 7.1 ´ 10–3, Ka2 = 6.2 ´ 10–8, and Ka3 = 4.5 ´ 10–13) A. 1.82 B. 2.12 C. 6.91 D. 7.21 E. 12.44
20. If the acid to base ratio in a buffer increases by a factor of 10, the pH of the buffer A. decreases by 10. B. decreases by 1. C. increases by 1. D. increases by 10. E. is not affected.
21. The buffer capacity is A. the amount of H+ or OH- that can be absorbed by a buffer without changing the pH appreciably. B. the amount of buffer required to neutralize a strong acid or base. C. a measure of the volume of buffer that a container may hold. D. the volume of buffer required to neutralize 1.00 L of 1.00 M HCl. E. dependent only on the ratio of base to acid in a buffer.
22. What is the effect on pH when a buffer is diluted by a factor of 10? A. The buffer pH deceases by 10. B. The buffer pH decreases by 1. C. The buffer pH decreases by 1. D. The buffer pH increases by 10. E. The buffer pH does not change appreciably.
23. The Ka of bicarbonate ion, HCO3–, is 4.8 ´ 10–11. What [CO32–]/[HCO3–] ratio is necessary to make a buffer with a pH of 11.00? A. 0.21 B. 0.32 C. 0.68 D. 4.8 E. 6.8
24. The Ka for H2PO4- is 6.2 ´ 10-8. What volume of 0.10 M NaOH must be added to 0.50 L of 0.30 M H2PO4- to make a pH = 7.50 buffer? A. 0.10 L B. 0.13 L C. 0.25 L D. 0.51 L E. 0.99 L
25. A buffer is prepared by adding 1.50 ´ 102 mL of 0.250 M NaOH to 2.50 ´ 102 mL of 0.350 M weak acid, HB. The solution is diluted to 1.00 L. If the pH of the resulting solution is 6.55, what is the pKa of the weak acid? A. 6.55 B. 6.67 C. 6.87 D. 6.92 E. 7.20
26. A buffer is composed of 0.400 mol H2PO4– and 0.400 mol HPO42– diluted with water to a volume of 1.00 L. The pH of the buffer is 7.210. How many moles of HCl must be added to decrease the pH to 6.210? A. 0.200 mol B. 0.327 mol C. 0.360 mol D. 0.400 mol E. 3.60 mol
27. A buffer contains 0.50 mol NH4+ and 0.50 mol NH3 diluted with water to 1.0 L. How many moles of NaOH are required to increase the pH of the buffer to 10.00? (pKa of NH4+ = 9.25) A. 0.035 mol B. 0.15 mol C. 0.35 mol D. 0.41 mol E. 2.8 mol
28. All of the following solutions are buffers EXCEPT A. 100 mL of 0.500 M H2PO4- added to 100 mL of 0.200 M HPO42-. B. 100 mL of 0.200 M HF added to 200 mL of 0.200 M NaF. C. 200 mL of 0.200 M HCl added to 200 mL of 0.400 M CH3CO2-. D. 200 mL of 0.500 M NaOH added to 200 mL of 1.000 M HF. E. 200 mL of 0.400 M HCl added to 100 mL of 0.200 M CH3CO2-.
29. Which of the following solutions has the greatest buffer capacity? A. A 1.0 L aqueous solution of 1.00 M HCl B. A 1.0 L mixture of 1.00 M H3PO4 and 1.00 M H2PO4C. A 1.0 L mixture of 1.00 M H3PO4 and 0.500 M NaOH D. A 1.0 L mixture of 1.00 M H2PO4- and 1.00 M HCl E. A 1.0 L aqueous solution of 0.75 M H2SO4
30. Which of the following solutions has the least buffer capacity? A. A 1.0 L mixture of 0.100 M H3PO4 and 0.100 M H2PO4B. A 1.0 L mixture of 0.200 M H3PO4 and 0.100 M NaOH C. A 1.0 L mixture of 0.200 M H2PO4- and 0.100 M HCl D. A 1.0 L mixture of 0.200 M H2PO4- and 0.100 M HNO3 E. All of the above have the same buffer capacity.
31. Which of the following conjugate acid-base pairs helps maintain blood at a pH of 7.40? A. H3O+ and H2O B. H2PO4- and HPO42C. H3PO4 and H2PO4D. H2CO3 and HCO3E. HCl and Cl-
32. Which of the following statements concerning acid-base indicators are true? 1. 2. 3.
Acid-base indicators are derived from weak acids. The acid and base forms of the indicator are different colors. Only small quantities of indicator are added to a solution.
A. 1 only B. 2 only C. 3 only D. 2 and 3 E. 1, 2, and 3 33. In a titration of HCl(aq) with NaOH(aq), the equivalence point is the point at which A. equal volumes of HCl and NaOH have been combined. B. the concentration of H+ equal 1.0 ´ 10-14 M. C. equal moles of HCl and NaOH have reacted. D. half of the HCl has been neutralized by hydroxide ion. E. the pH equals 14.00.
34. Which of the following statements concerning phenolphthalein are true? 1. 2. 3.
Phenolphthalein changes from colorless in acidic solutions to pink in basic solutions. Phenolphthalein is obtained from red cabbage juice. Phenolphthalein changes color at about pH 9.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 35. What is the relationship between the pKa for an acid-base indicator and the pH at which the indicator changes color? A. pH = Ka B. pH = pKa C. pH = -log[pKa] D. pH = -log[H+] E. [H+] = pKa
36. Which one of the following conditions is always true for a titration of a weak acid with a strong base? A. A colored indicator with a pKa less than 7 should be used. B. If a colored indicator is used, it must change color rapidly in the weak acid's buffer region. C. Equal volumes of weak acid and strong base are required to reach the equivalence point. D. The equivalence point occurs at a pH equal to 7. E. The equivalence point occurs at a pH greater than 7.
37. All of the following statements are false for the titration of a strong base by a strong acid EXCEPT A. the equivalence point occurs at pH 7. B. there will be at least two equivalence points, one for the base and one for the acid. C. the equivalence point and the end point are identical, regardless of what indicator is used. D. equal masses of acid and base are required to neutralize the solution. E. phenolphthalein, if used as an indicator, will turn from colorless to pink.
38. A 30.00 mL sample of vinegar is titrated with 0.4190 M NaOH(aq). If the titration requires 27.83 mL of NaOH(aq), what is the concentration of acetic acid in the vinegar? A. 0.1931 M B. 0.2016 M C. 0.2174 M D. 0.3887 M E. 0.4517 M
39. A 25.0 mL sample of 0.400 M NH3(aq) is titrated with 0.400 M HCl(aq). What is the pH at the equivalence point? (Kb of NH3 = 1.8 ´ 10–5) A. 2.72 B. 4.97 C. 7.00 D. 9.03 E. 11.28
40. A volume of 25.0 mL of 0.100 M C6H5CO2H(aq) is titrated with 0.100 M NaOH(aq). What is the pH after the addition of 12.5 mL of NaOH? (Ka of benzoic acid = 6.3 ´ 10–5) A. 2.60 B. 4.20 C. 5.40 D. 7.00 E. 8.60
41. If 25.00 mL of 0.200 M HCl is titrated with 0.200 M NaOH, what is the pH of the solution after the addition of 22.50 mL of NaOH? A. 0.58 B. 1.98 C. 4.67 D. 6.88 E. 10.70
42. Potassium hydrogen phthalate (KHP) is used to standardize sodium hydroxide. If 35.39 mL of NaOH(aq) is required to titrate 0.8246 g KHP to the equivalence point, what is the concentration of the NaOH(aq)? (The molar mass of KHP = 204.2 g/mol) HC8H4O4–(aq) + OH–(aq) C8H4O42–(aq) + H2O(l) A. 0.02318 M B. 0.05705 M C. 0.0859 M D. 0.1141 M E. 0.1429 M
43. What volume of 0.2045 M NaOH is necessary to titrate 50.00 mL of 0.1177 M acetic acid? A. 12.02 mL B. 18.01 mL C. 23.77 mL D. 28.78 mL E. 86.87 mL
44. If 25.00 mL of 0.2500 M formic acid is titrated with 25.00 mL of 0.2500 M NaOH, what is the pH at the equivalence point? The Ka for formic acid is 1.8 ´ 10-4. A. 6.79 B. 8.42 C. 9.01 D. 9.67 E. 10.04
45. Sodium carbonate, Na2CO3, can be titrated with HCl according to the following balanced chemical equation. 2H+(aq) + CO32-(aq) ® H2O(l) + CO2(g) What volume of 0.2000 M HCl is necessary to titrate 0.3462 g of sodium carbonate? A. 16.33 mL B. 21.74 mL C. 32.67 mL D. 37.79 mL E. 41.11 mL
46. An impure sample of sodium carbonate, Na2CO3, is titrated with 0.113 M HCl according to the reaction below. 2 HCl(aq) + Na2CO3(aq) CO2(g) + H2O(l) + 2 NaCl(aq) What is the percent of Na2CO3 in a 0.613 g sample if the titration requires 26.14 mL of HCl? The molar mass of Na2CO3 is 106.0 g/mol. A. 0.295% B. 15.7% C. 25.5% D. 51.1% E. 67.9%
47. Which indicator is most appropriate for the titration of acetic acid with NaOH? The Ka for CH3CO2H is 1.8 ´ 10-5. A. methyl red (pH 5) B. bromthymol blue (pH 7) C. phenolphthalein (pH 9) D. Both methyl red and bromthymol blue are suitable. E. All three indicators are suitable.
48. Which is the best colored indicator to use in the titration of 0.1 M CH3CO2H(aq) with NaOH(aq)? Why? (Ka of CH3CO2H = 1.8 ´ 10–5, Kb of CH3CO2– = 5.6 ´ 10–10)
Indicator Bromcresol Green Bromthymol Blue Phenolphthalein
pKa 4.8 6.8 9.2
A. Bromcresol Green. The equivalence point for a weak acid titration occurs at low pH. B. Bromthymol Blue. The pH at the equivalence point is 7.0. C. Bromcresol Green. The pKa of CH3CO2H and the pKa of the indicator are similar. D. Phenolphthalein. The pKa of CH3CO2– and the pKb of the indicator are similar. E. Phenolphthalein. The pH at the equivalence point is near the pKa of the indicator. 49. In a titration experiment, if the initial solution pH is 4.0 and the equivalence point occurs at pH 9.0, then the reaction corresponds to A. the titration of a weak acid by a strong base. B. the titration of a weak acid by a weak base. C. the titration of a weak base by a strong acid. D. the titration of a strong acid by a strong base. E. the titration of a strong base by a strong acid.
50. Which two acids are responsible for acid rain? A. HCl and HNO3 B. HCl and H2SO4 C. HNO3 and HF D. HClO4 and HNO3 E. HNO3 and H2SO4
51. Which one of the following reactions describes the attack of acid rain on limestone? A. 2H+(aq) + Ca(s) ® Ca2+(aq) + H2(g) B. 2H+(aq) + Ca(OH)2(s) ® Ca2+(aq) + 2H2O(l) C. 2H+(aq) + CaO(s) ® Ca2+(aq) + H2O(l) D. 2H+(aq) + CaCO3(s) ® Ca2+(aq) + CO2(g) + H2O(l) E. 3H+(aq) + Al(OH)3(s) ® Al3+(aq) + 3H2O(l)
52. Natural rainfall is slightly acidic at a pH of about 5.5. Why is the pH of natural rain acidic? A. pure water is acidic B. HNO3 forms when N2 dissolves in rainwater C. H2CO3 forms when CO2 dissolves in rainwater D. H2O3 forms when O2 dissolves in rainwater E. H2SO4 forms when smog dissolves in rainwater
53. A 20.00-mL sample of 0.220 M triethylamine, (CH3CH2)3N is titrated with 0.544 M HCl (Kb(CH3CH2)3N = 5.2 ´ 10-4). How many mL of HCl are required to reach the equivalence point? A. 16.17 mL B. 21.74 mL C. 32.67 mL D. 8.09 mL E. 49.45 mL
54. A 20.00-mL sample of 0.220 M triethylamine, (CH3CH2)3N is titrated with 0.544 M HCl (Kb(CH3CH2)3N = 5.2 ´ 10-4). What is the pH of the solution at the equivalence point? A. 5.76 B. 8.24 C. 7.00 D. 3.28 E. 10.72
55. Which of the following acids build up in muscles that are overexerted, causing pain? A. Hydrochloric acid B. Acetic acid C. Carbonic acid D. Hypochlorous acid E. Lactic acid
Chapter 14--Equilibria in Acid-Base Solutions Key
1. An acid-base equilibrium system is created by dissolving 0.50 mol CH3CO2H in water to a volume of 1.0 L. What is the effect of adding 0.50 mol CH3CO2–(aq) to this solution? 1. 2. 3.
The pH of the solution will equal 7.00 because equal concentrations of a weak acid and its conjugate base are present. Some CH3CO2H(aq) will ionize, increasing the concentration of CH3CO2–(aq) and increasing the pH. Some CH3CO2–(aq) will react with H3O+(aq), increasing the concentration of CH3CO2H(aq) and reestablishing the solution equilibrium.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 2. What is the effect of adding NaOH(aq) to an aqueous solution of ammonia? 1. 2. 3.
The pH of the solution will increase. The concentration of NH4+(aq) will decrease. The concentration of NH3(aq) will increase.
A. 1 only B. 2 only C. 3 only D. 2 and 3 E. 1, 2, and 3 3. All of the following statements concerning acid-base buffers are true EXCEPT A. buffers are resistant to pH changes upon addition of small quantities of strong acids or bases. B. buffers are used as colored indicators in acid-base titrations. C. the pH of a buffer is close to the pKa of the weak acid from which it is made. D. buffers contain appreciable quantities of a weak acid and its conjugate base. E. buffers are resistant to changes in pH when diluted with water.
4. Which of the following pairs will form a buffer when mixed together in an aqueous solution? A. KCl and KH2PO4 B. HCl and KOH C. Ca(OH)2 and NaOH D. HF and NaF E. None of the above will form a buffer.
5. When mixed together, all of the following pairs can form buffers EXCEPT A. H3PO4 and NaH2PO4. B. NaH2PO4 and Na2HPO4. C. CH3CO2H and NaOH. D. HCl and NaCH3CO2. E. NaI and NaOH.
6. Which of the following equations is the Henderson-Hasselbalch equation?
A.
B.
C.
D.
E.
7. What is the pH of a solution that results from diluting 0.30 mol acetic acid (CH3CO2H) and 0.20 mol sodium acetate (NaCH3CO2) with water to a volume of 1.0 L? (Ka of CH3CO2H = 1.8 ´ 10–5) A. 4.35 B. 4.57 C. 4.74 D. 4.92 E. 5.14
8. What is the pH of an aqueous solution composed of 0.64 M NH4+ and 0.20 M NH3? (Ka of NH4+ = 5.6 ´ 10–10) A. 4.80 B. 8.75 C. 9.20 D. 9.25 E. 9.76
9. The Ka of hypobromous acid, HOBr, is 2.6 ´ 10-9. Calculate the pH of a solution which is composed of 0.40 M HOBr and 0.40 M NaOBr. A. 0.40 B. 0.80 C. 4.49 D. 8.59 E. 9.12
10. The Ka of acetic acid is 1.8 ´ 10-5. Calculate the pH of a solution that is composed of 1.0 M CH3CO2H and 0.50 M NaCH3CO2. A. 2.24 B. 4.44 C. 4.74 D. 5.05 E. 6.99
11. A buffer may be prepared by mixing a weak acid with a roughly equivalent amount of strong base. Which of the acids below is best for the preparation of a buffer with a pH of 4.00? A. sulfurous acid , H2SO3; Ka = 1.7 ´ 10-2 B. hydrofluoric acid, HF; Ka = 6.9 ´ 10-4 C. benzoic acid, HC7H5O2; Ka = 6.6 ´ 10-5 D. dihydrogen phosphate ion, H2PO4-; Ka = 6. 2 ´ 10-8 E. ammonium ion, NH4+; Ka = 5.6 ´ 10-10
12. A buffer may be prepared by mixing a weak acid with a roughly equivalent amount of strong base. Which of the acids below is best for the preparation of a buffer with a pH of 9.00? A. chlorous acid, HClO2; Ka = 2.8 ´ 10-8 B. formic acid, HCO2H; Ka = 1.9 ´ 10-4 C. benzoic acid, HC7H5O2; Ka = 6.6 ´ 10-5 D. dihydrogen phosphate ion, H2PO4-; Ka = 6.2 ´ 10-8 E. ammonium ion, NH4+; Ka = 5.6 ´ 10-10
13. What is the pH of the buffer that results when 12.0 g of NaH2PO4 and 8.00 g of Na2HPO4 are diluted with water to a volume of 0.50 L? (Ka of H2PO4– = 6.2 ´ 10–8, the molar masses of NaH2PO4 and Na2HPO4 are 120.0 g/mol and 142.0 mol, respectively) A. 4.10 B. 6.96 C. 7.21 D. 7.46 E. 9.90
14. An acetic acid-sodium acetate buffer is prepared by mixing 6.00 g NaCH3CO2 with 4.00 mL of CH3CO2H (density = 1.042 g/mL) and diluting to 1.00 L. What is the pH of the resulting buffer? The Ka of acetic acid is 1.8 ´ 10-5. A. 4.10 B. 4.77 C. 4.99 D. 5.19 E. 5.78
15. What is the pH of the buffer that results when 32 g sodium acetate (NaCH3CO2) is mixed with 500.0 mL of 1.0 M acetic acid (CH3CO2H) and diluted with water to 1.0 L? (Ka of CH3CO2H = 1.8 ´ 10–5) A. 2.52 B. 4.23 C. 4.44 D. 4.64 E. 4.74
16. A buffer is prepared by combining 25 mL of 0.50 M NaF(aq) with 25 mL of 0.25 M HCl. What is the pH of the buffer? (Ka (HF) = 6.9 ´ 10-4) A. 3.16 B. 3.50 C. 4.12 D. 4.60 E. 7.11
17. A buffer is prepared by combining 25 mL of 0.50 M NH3(aq) with 25 mL of 0.20 M HCl. What is the pH of the buffer? (Ka (NH4+) = 5.6 ´ 10-10) A. 7.76 B. 8.00 C. 8.86 D. 9.43 E. 9.65
18. A buffer is prepared by combining 10.0 g of NaH2PO4 with 150 mL of 0.20 M NaOH and diluting to 2.0 L. What is the pH of the buffer? (Ka (H2PO4-) = 6.2 ´ 10-8) A. 5.12 B. 6.76 C. 6.96 D. 7.20 E. 7.45
19. What is the pH of a buffer that results when 0.50 mole of H3PO4 is mixed with 0.75 mole of NaOH and diluted with water to 1.00 L? (The acid dissociation constants of phosphoric acid are Ka1 = 7.1 ´ 10–3, Ka2 = 6.2 ´ 10–8, and Ka3 = 4.5 ´ 10–13) A. 1.82 B. 2.12 C. 6.91 D. 7.21 E. 12.44
20. If the acid to base ratio in a buffer increases by a factor of 10, the pH of the buffer A. decreases by 10. B. decreases by 1. C. increases by 1. D. increases by 10. E. is not affected.
21. The buffer capacity is A. the amount of H+ or OH- that can be absorbed by a buffer without changing the pH appreciably. B. the amount of buffer required to neutralize a strong acid or base. C. a measure of the volume of buffer that a container may hold. D. the volume of buffer required to neutralize 1.00 L of 1.00 M HCl. E. dependent only on the ratio of base to acid in a buffer.
22. What is the effect on pH when a buffer is diluted by a factor of 10? A. The buffer pH deceases by 10. B. The buffer pH decreases by 1. C. The buffer pH decreases by 1. D. The buffer pH increases by 10. E. The buffer pH does not change appreciably.
23. The Ka of bicarbonate ion, HCO3–, is 4.8 ´ 10–11. What [CO32–]/[HCO3–] ratio is necessary to make a buffer with a pH of 11.00? A. 0.21 B. 0.32 C. 0.68 D. 4.8 E. 6.8
24. The Ka for H2PO4- is 6.2 ´ 10-8. What volume of 0.10 M NaOH must be added to 0.50 L of 0.30 M H2PO4- to make a pH = 7.50 buffer? A. 0.10 L B. 0.13 L C. 0.25 L D. 0.51 L E. 0.99 L
25. A buffer is prepared by adding 1.50 ´ 102 mL of 0.250 M NaOH to 2.50 ´ 102 mL of 0.350 M weak acid, HB. The solution is diluted to 1.00 L. If the pH of the resulting solution is 6.55, what is the pKa of the weak acid? A. 6.55 B. 6.67 C. 6.87 D. 6.92 E. 7.20
26. A buffer is composed of 0.400 mol H2PO4– and 0.400 mol HPO42– diluted with water to a volume of 1.00 L. The pH of the buffer is 7.210. How many moles of HCl must be added to decrease the pH to 6.210? A. 0.200 mol B. 0.327 mol C. 0.360 mol D. 0.400 mol E. 3.60 mol
27. A buffer contains 0.50 mol NH4+ and 0.50 mol NH3 diluted with water to 1.0 L. How many moles of NaOH are required to increase the pH of the buffer to 10.00? (pKa of NH4+ = 9.25) A. 0.035 mol B. 0.15 mol C. 0.35 mol D. 0.41 mol E. 2.8 mol
28. All of the following solutions are buffers EXCEPT A. 100 mL of 0.500 M H2PO4- added to 100 mL of 0.200 M HPO42-. B. 100 mL of 0.200 M HF added to 200 mL of 0.200 M NaF. C. 200 mL of 0.200 M HCl added to 200 mL of 0.400 M CH3CO2-. D. 200 mL of 0.500 M NaOH added to 200 mL of 1.000 M HF. E. 200 mL of 0.400 M HCl added to 100 mL of 0.200 M CH3CO2-.
29. Which of the following solutions has the greatest buffer capacity? A. A 1.0 L aqueous solution of 1.00 M HCl B. A 1.0 L mixture of 1.00 M H3PO4 and 1.00 M H2PO4C. A 1.0 L mixture of 1.00 M H3PO4 and 0.500 M NaOH D. A 1.0 L mixture of 1.00 M H2PO4- and 1.00 M HCl E. A 1.0 L aqueous solution of 0.75 M H2SO4
30. Which of the following solutions has the least buffer capacity? A. A 1.0 L mixture of 0.100 M H3PO4 and 0.100 M H2PO4B. A 1.0 L mixture of 0.200 M H3PO4 and 0.100 M NaOH C. A 1.0 L mixture of 0.200 M H2PO4- and 0.100 M HCl D. A 1.0 L mixture of 0.200 M H2PO4- and 0.100 M HNO3 E. All of the above have the same buffer capacity.
31. Which of the following conjugate acid-base pairs helps maintain blood at a pH of 7.40? A. H3O+ and H2O B. H2PO4- and HPO42C. H3PO4 and H2PO4D. H2CO3 and HCO3E. HCl and Cl-
32. Which of the following statements concerning acid-base indicators are true? 1. 2. 3.
Acid-base indicators are derived from weak acids. The acid and base forms of the indicator are different colors. Only small quantities of indicator are added to a solution.
A. 1 only B. 2 only C. 3 only D. 2 and 3 E. 1, 2, and 3
33. In a titration of HCl(aq) with NaOH(aq), the equivalence point is the point at which A. equal volumes of HCl and NaOH have been combined. B. the concentration of H+ equal 1.0 ´ 10-14 M. C. equal moles of HCl and NaOH have reacted. D. half of the HCl has been neutralized by hydroxide ion. E. the pH equals 14.00.
34. Which of the following statements concerning phenolphthalein are true? 1. 2. 3.
Phenolphthalein changes from colorless in acidic solutions to pink in basic solutions. Phenolphthalein is obtained from red cabbage juice. Phenolphthalein changes color at about pH 9.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 35. What is the relationship between the pKa for an acid-base indicator and the pH at which the indicator changes color? A. pH = Ka B. pH = pKa C. pH = -log[pKa] D. pH = -log[H+] E. [H+] = pKa
36. Which one of the following conditions is always true for a titration of a weak acid with a strong base? A. A colored indicator with a pKa less than 7 should be used. B. If a colored indicator is used, it must change color rapidly in the weak acid's buffer region. C. Equal volumes of weak acid and strong base are required to reach the equivalence point. D. The equivalence point occurs at a pH equal to 7. E. The equivalence point occurs at a pH greater than 7.
37. All of the following statements are false for the titration of a strong base by a strong acid EXCEPT A. the equivalence point occurs at pH 7. B. there will be at least two equivalence points, one for the base and one for the acid. C. the equivalence point and the end point are identical, regardless of what indicator is used. D. equal masses of acid and base are required to neutralize the solution. E. phenolphthalein, if used as an indicator, will turn from colorless to pink.
38. A 30.00 mL sample of vinegar is titrated with 0.4190 M NaOH(aq). If the titration requires 27.83 mL of NaOH(aq), what is the concentration of acetic acid in the vinegar? A. 0.1931 M B. 0.2016 M C. 0.2174 M D. 0.3887 M E. 0.4517 M
39. A 25.0 mL sample of 0.400 M NH3(aq) is titrated with 0.400 M HCl(aq). What is the pH at the equivalence point? (Kb of NH3 = 1.8 ´ 10–5) A. 2.72 B. 4.97 C. 7.00 D. 9.03 E. 11.28
40. A volume of 25.0 mL of 0.100 M C6H5CO2H(aq) is titrated with 0.100 M NaOH(aq). What is the pH after the addition of 12.5 mL of NaOH? (Ka of benzoic acid = 6.3 ´ 10–5) A. 2.60 B. 4.20 C. 5.40 D. 7.00 E. 8.60
41. If 25.00 mL of 0.200 M HCl is titrated with 0.200 M NaOH, what is the pH of the solution after the addition of 22.50 mL of NaOH? A. 0.58 B. 1.98 C. 4.67 D. 6.88 E. 10.70
42. Potassium hydrogen phthalate (KHP) is used to standardize sodium hydroxide. If 35.39 mL of NaOH(aq) is required to titrate 0.8246 g KHP to the equivalence point, what is the concentration of the NaOH(aq)? (The molar mass of KHP = 204.2 g/mol) HC8H4O4–(aq) + OH–(aq) C8H4O42–(aq) + H2O(l) A. 0.02318 M B. 0.05705 M C. 0.0859 M D. 0.1141 M E. 0.1429 M
43. What volume of 0.2045 M NaOH is necessary to titrate 50.00 mL of 0.1177 M acetic acid? A. 12.02 mL B. 18.01 mL C. 23.77 mL D. 28.78 mL E. 86.87 mL
44. If 25.00 mL of 0.2500 M formic acid is titrated with 25.00 mL of 0.2500 M NaOH, what is the pH at the equivalence point? The Ka for formic acid is 1.8 ´ 10-4. A. 6.79 B. 8.42 C. 9.01 D. 9.67 E. 10.04
45. Sodium carbonate, Na2CO3, can be titrated with HCl according to the following balanced chemical equation. 2H+(aq) + CO32-(aq) ® H2O(l) + CO2(g) What volume of 0.2000 M HCl is necessary to titrate 0.3462 g of sodium carbonate? A. 16.33 mL B. 21.74 mL C. 32.67 mL D. 37.79 mL E. 41.11 mL
46. An impure sample of sodium carbonate, Na2CO3, is titrated with 0.113 M HCl according to the reaction below. 2 HCl(aq) + Na2CO3(aq) CO2(g) + H2O(l) + 2 NaCl(aq) What is the percent of Na2CO3 in a 0.613 g sample if the titration requires 26.14 mL of HCl? The molar mass of Na2CO3 is 106.0 g/mol. A. 0.295% B. 15.7% C. 25.5% D. 51.1% E. 67.9%
47. Which indicator is most appropriate for the titration of acetic acid with NaOH? The Ka for CH3CO2H is 1.8 ´ 10-5. A. methyl red (pH 5) B. bromthymol blue (pH 7) C. phenolphthalein (pH 9) D. Both methyl red and bromthymol blue are suitable. E. All three indicators are suitable.
48. Which is the best colored indicator to use in the titration of 0.1 M CH3CO2H(aq) with NaOH(aq)? Why? (Ka of CH3CO2H = 1.8 ´ 10–5, Kb of CH3CO2– = 5.6 ´ 10–10)
Indicator Bromcresol Green Bromthymol Blue Phenolphthalein
pKa 4.8 6.8 9.2
A. Bromcresol Green. The equivalence point for a weak acid titration occurs at low pH. B. Bromthymol Blue. The pH at the equivalence point is 7.0. C. Bromcresol Green. The pKa of CH3CO2H and the pKa of the indicator are similar. D. Phenolphthalein. The pKa of CH3CO2– and the pKb of the indicator are similar. E. Phenolphthalein. The pH at the equivalence point is near the pKa of the indicator. 49. In a titration experiment, if the initial solution pH is 4.0 and the equivalence point occurs at pH 9.0, then the reaction corresponds to A. the titration of a weak acid by a strong base. B. the titration of a weak acid by a weak base. C. the titration of a weak base by a strong acid. D. the titration of a strong acid by a strong base. E. the titration of a strong base by a strong acid.
50. Which two acids are responsible for acid rain? A. HCl and HNO3 B. HCl and H2SO4 C. HNO3 and HF D. HClO4 and HNO3 E. HNO3 and H2SO4
51. Which one of the following reactions describes the attack of acid rain on limestone? A. 2H+(aq) + Ca(s) ® Ca2+(aq) + H2(g) B. 2H+(aq) + Ca(OH)2(s) ® Ca2+(aq) + 2H2O(l) C. 2H+(aq) + CaO(s) ® Ca2+(aq) + H2O(l) D. 2H+(aq) + CaCO3(s) ® Ca2+(aq) + CO2(g) + H2O(l) E. 3H+(aq) + Al(OH)3(s) ® Al3+(aq) + 3H2O(l)
52. Natural rainfall is slightly acidic at a pH of about 5.5. Why is the pH of natural rain acidic? A. pure water is acidic B. HNO3 forms when N2 dissolves in rainwater C. H2CO3 forms when CO2 dissolves in rainwater D. H2O3 forms when O2 dissolves in rainwater E. H2SO4 forms when smog dissolves in rainwater
53. A 20.00-mL sample of 0.220 M triethylamine, (CH3CH2)3N is titrated with 0.544 M HCl (Kb(CH3CH2)3N = 5.2 ´ 10-4). How many mL of HCl are required to reach the equivalence point? A. 16.17 mL B. 21.74 mL C. 32.67 mL D. 8.09 mL E. 49.45 mL
54. A 20.00-mL sample of 0.220 M triethylamine, (CH3CH2)3N is titrated with 0.544 M HCl (Kb(CH3CH2)3N = 5.2 ´ 10-4). What is the pH of the solution at the equivalence point? A. 5.76 B. 8.24 C. 7.00 D. 3.28 E. 10.72
55. Which of the following acids build up in muscles that are overexerted, causing pain? A. Hydrochloric acid B. Acetic acid C. Carbonic acid D. Hypochlorous acid E. Lactic acid
Chapter 15--Complex Ion and Precipitation Equilibria 1. Which of the following equations represents the formation constant of Ti(NH3)63+? A. Ti3+(aq) + (NH3)6(aq) Ti(NH3)63+(aq) 3+ 3+ B. Ti(NH3)6 (aq) Ti (aq) + 6NH3(aq) 3+ C. Ti (aq) + NH3(aq) Ti(NH3)3+(aq) 3+ D. 6Ti (aq) + 6NH3(aq) Ti(NH3)63+(aq) 3+ E. Ti (aq) + 6NH3(aq) Ti(NH3)63+(aq)
2. The formation constant for Zn(NH3)42+ is 3.6 ´ 108. What is the ratio of Zn(NH3)42+ to Zn2+ in 0.010 M NH3? A. 3.6 B. 1.4 ´ 103 C. 3.6 ´ 106 D. 1.4 ´ 103 E. 3.6 ´ 1010
3. The formation constant for Ag(CN)2-(aq) is 2 ´ 1020. At what cyanide ion concentration is [Ag+] = [ Ag(CN)2-]? A. 3 ´ 10-41 M B. 5 ´ 10-21 M C. 7 ´ 10-11 M D. 6 ´ 10-6 M E. 1 ´ 1010 M
4. Ksp stands for A. saturated precipitate. B. soluble precipitate. C. soluble particle. D. saturated precipitate. E. solubility product.
5. What is the chemical equation that describes the dissolution of calcium iodate in water? A. CaI2(s) Ca2+(aq) + 2I-(aq) B. Ca2+(aq) + IO3-(aq) CaIO3(s) C. Ca(IO3)2(s) Ca(s) + I2(aq) + 3O2(g) D. Ca(IO3)2(s) Ca2+(aq) + 2IO3-(aq) E. Ca(IO3)2(s) CaI2(aq) + 3O2(g)
6. What is the chemical equation that describes the dissolution of nickel(II) hydroxide in water? A. Ni(OH)2(s) Ni2+(aq) + 2OH-(aq) 2+ B. Ni (aq) + 3OH (aq) Ni(OH)2(s) + C. Ni2OH(s) + H (aq) 2Ni+(aq) + H2O(l) + D. Ni(OH)2(s) + 2H (aq) Ni2+(aq) + 2H2O(l) E. Ni(OH)2(aq) + H2O(l) Ni(OH)3(s) + H+(aq)
7. What is the chemical equation that describes the dissolution of silver oxalate, Ag2(C2O4), in water? A. Ag2(C2O4)(s) Ag22+(aq) + C2O42-(aq) B. Ag2(C2O4)(s) 2Ag+(aq) + C2O42-(aq) C. Ag2(C2O4)(s) Ag2+(aq) + C2O4-(aq) + 2D. 2Ag (aq) + C2O4 (aq) Ag2(C2O4)(s) + E. Ag2 (aq) + C2O4 (aq) Ag2(C2O4)(s)
8. Which of the following equations is the solubility product of Cu(II) hydroxide, Cu(OH)2?
A.
B. C. D. E.
9. What is the equilibrium constant expression for the dissolution of solid lead oxalate, PbC2O4, in water? A. B. Ksp = [Pb]2 [C2O4]2
C.
D. E.
10. What is the equilibrium constant expression for the dissolution of solid calcium fluoride in water? A. Ksp = [Ca2+][F-] B. Ksp = [Ca2+]2[F-] C. Ksp = [Ca2+][F-]2
D.
E.
11. Consider the equilibrium of PbCl2(s) in water. PbCl2(s) Pb2+(aq) + 2Cl-(aq) What is the effect of adding NaCl(aq) to the equilibrium solution? A. The sodium ion reduces the Pb2+ to Pb(s). B. PbCl2 solubility increases due to the common-ion effect. C. PbNa2(s) precipitates. D. The NaCl(aq) has no effect on the system. E. PbCl2(s) precipitates until equilibrium is reestablished.
12. Consider the equilibrium of Ca(OH)2(s) in water. Ca(OH)2(s) Ca2+(aq) + 2OH-(aq) What is the effect of raising the pH of the solution? A. Ca2+(aq) is reduced to Ca(s). B. The concentration of hydronium ion increases. C. The concentration of Ca2+ increases as Ca(OH)2 dissolves. D. Ca(OH)2(s) precipitates until equilibrium is reestablished. E. Hydroxide ion is reduced to H2(g).
13. Consider the equilibrium of lead sulfide, PbS, in water. PbS(s) Pb2+(aq) + S2-(aq) What is the effect of adding PbS(s) to the solution? A. The PbS(s) has no effect on the system. B. More Pb2+ and S2- will form. C. More Pb2+ will form, but the concentration of S2- will remain constant. D. More S2- will form, but the concentration of Pb2+ will remain constant. E. The concentrations of Pb2+ and S2- will decrease.
14. Consider the equilibrium of magnesium fluoride, MgF2(s), in water. MgF2(s) Mg2+(aq) + 2F-(aq) What is the effect of adding perchloric acid (i.e. decreasing the pH of solution)? A. The acid will have no effect on the magnesium fluoride equilibrium. B. The F- will be protonated, resulting in more MgF2(s) dissolving. C. More MgF2(s) will be formed, thus reducing Mg2+ and F- concentrations in solution. D. Magnesium perchlorate, Mg(ClO4)2(s), will precipitate. E. Magnesium hydride, MgH2(s), will precipitate.
15. The Ksp of barium chromate is 1.2 ´ 10-10. What is the concentration of Ba2+ in equilibrium with BaCrO4(s) if [CrO42-] = 4.3 ´ 10-3 M? A. 5.1 ´ 10-13 M B. 2.8 ´ 10-8 M C. 1.1 ´ 10-5 M D. 1.7 ´ 10-4 M E. 2.5 ´ 10-3 M
16. The Ksp of calcium iodate is 7.1 ´ 10-7. What is the concentration of Ca2+ in equilibrium with Ca(IO3)2(s) if [IO3-] = 3.3 ´ 10-2 M? A. 3.3 ´ 10-7 M B. 2.2 ´ 10-5 M C. 8.1 ´ 10-5 M D. 6.5 ´ 10-4 M E. 7.1 ´ 10-3 M
17. The Ksp of aluminum hydroxide, Al(OH)3, is 2 ´ 10-31. What pH is required to limit the Al3+ concentration to less than or equal to 1 ´ 10-10 M? A. 3.6 B. 6.4 C. 7.1 D. 7.8 E. 11.5
18. The Ksp of calcium phosphate is 1 ´ 10-33. What is the concentration of Ca2+ in equilibrium with Ca3(PO4)2(s) if [PO43-] = 1 ´ 10-5 M? A. 1 ´ 10-28 M B. 1 ´ 10-23 M C. 1 ´ 10-11 M D. 5 ´ 10-10 M E. 2 ´ 10-8 M
19. The Ksp of Fe(OH)3(s) is 3 ´ 10-39. What concentration of Fe3+ can exist in solution at pH 3.0? A. 3 ´ 10-4 M B. 3 ´ 10-6 M C. 3 ´ 10-21 M D. 3 ´ 10-24 M E. 3 ´ 10-30 M
20. The Ksp of BaSO4 is 1.1 ´ 10–10 at 25C. What mass of BaSO4 (molar mass = 233.4 g/mol) will dissolve in 1.0 L of water at 25C? A. 2.6 ´ 10-8 g B. 4.5 ´ 10-8 g C. 1.0 ´ 10-5 g D. 1.6 ´ 10-4 g E. 2.4 ´ 10-3 g
21. For AgCl, Ksp = 1.8 ´ 10-10. What will occur if 250 mL of 1.5 ´ 10-3 M NaCl is mixed with 250 mL of 2.0 ´ 10-7 M AgNO3? A. Q > Ksp. A precipitate will form. B. Ksp > Q. A precipitate will form. C. Q = Ksp. No precipitate will form. D. Q > Ksp. No precipitate will form. E. Ksp > Q. No precipitate will form.
22. For PbCl2, Ksp = 1.7 ´ 10-5. What will occur if 250 mL of 0.12 M Pb(NO3)2 is mixed with 250 mL of 0.070 M NaCl? A. Q > Ksp. A precipitate will form. B. Ksp > Q. A precipitate will form. C. Q = Ksp. No precipitate will form. D. Q > Ksp. No precipitate will form. E. Ksp > Q. No precipitate will form.
23. For Ca(OH)2, Ksp = 4.0 ´ 10-6. What will occur if 1.0 L of 0.100 M Ca(NO3)2 is prepared in a solution that is buffered at pH 12.50? A. Q > Ksp. A precipitate will form. B. Ksp > Q. A precipitate will form. C. Q = Ksp. No precipitate will form. D. Q > Ksp. No precipitate will form. E. Ksp > Q. No precipitate will form.
24. For Mg(OH)2, Ksp = 6 ´ 10-12. What will occur if 1.0 L of 0.010 M Mg(NO3)2 is prepared at pH 10.00? A. Q > Ksp. A precipitate will form. B. Ksp > Q. A precipitate will form. C. Q = Ksp. No precipitate will form. D. Q > Ksp. No precipitate will form. E. Ksp > Q. No precipitate will form.
25. What is the water solubility of AgI (Ksp = 1 ´ 10-16, MM = 234.8 g/mol) in moles per liter? A. 1 ´ 10-32 M B. 1 ´ 10-16 M C. 1 ´ 10-8 M D. 2 ´ 10-6 M E. 4 ´ 10-11 M
26. What is the water solubility of PbI2 (Ksp = 8.4 ´ 10-9, MM = 461 g/mol) in moles per liter? A. 8.4 ´ 10-9 M B. 1.9 ´ 10-7 M C. 9.1 ´ 10-5 M D. 1.3 ´ 10-3 M E. 2.0 ´ 10-3 M
27. What is the water solubility of Hg2Br2 (Ksp = 6 ´ 10-23) in moles per liter? A. 8 ´ 10-12 M B. 1 ´ 10-8 M C. 2 ´ 10-8 M D. 4 ´ 10-8 M E. 3 ´ 10-6 M
28. What is the water solubility of AgCl (Ksp = 1.8 ´ 10-10, MM = 143.4 g/mol) in grams per liter? A. 1.3 ´ 10-12 g/L B. 2.6 ´ 10-8 g/L C. 9.4 ´ 10-8 g/L D. 9.6 ´ 10-4 g/L E. 1.9 ´ 10-3 g/L
29. The Ksp of BaSO4 is 1.1 ´ 10–10 at 25C. What mass of BaSO4 (molar mass = 233.4 g/mol) will dissolve in 1.0 L of water at 25C? A. 2.6 ´ 10-8 g B. 4.5 ´ 10-8 g C. 1.0 ´ 10-5 g D. 1.6 ´ 10-4 g E. 2.4 ´ 10-3 g
30. What volume of water is needed to completely dissolve 100 g of AgBr (Ksp = 5 ´ 10-13, 187.8 g/mol)? A. 1 ´ 10-2 L B. 8 ´ 105 L C. 1 ´ 108 L D. 3 ´ 1010 L E. 1 ´ 1012 L
31. What volume of water is needed to completely dissolve 1.0 g of Ag2CrO4 (Ksp = 8.0 ´ 10-12, MM = 331.8 g/mol)? A. 0.94 L B. 1.1 L C. 1.7 L D. 24 L E. 2.6 ´ 105 L
32. The Ksp of Ca(OH)2 is 5.5 ´ 10–5 at 25C. What is the concentration of OH–(aq) in a saturated solution of Ca(OH)2(aq)? A. 1.9 ´ 10–3 M B. 7.4 ´ 10–3 M C. 2.4 ´ 10–2 M D. 4.0 ´ 10–2 M E. 4.8 ´ 10–2 M
33. A saturated solution of lead(II) sulfate can be prepared by diluting 0.0101 g of PbSO4 to 250 mL. What is the Ksp of lead(II) sulfate? A. 8.3 ´ 10-12 B. 1.6 ´ 10-9 C. 1.8 ´ 10-8 D. 4.5 ´ 10-8 E. 8.4 ´ 10-7
34. At pH 10.0, only 0.019 g of MgCl2 will dissolve per 1 L of solution. What is the Ksp of magnesium hydroxide, Mg(OH)2? A. 2 ´ 10-13 B. 2 ´ 10-12 C. 2 ´ 10-10 D. 2 ´ 10-8 E. 2 ´ 10-6
35. The solubility of BaCrO4(s) in water is 3.7 milligrams in 1.0 L at 25C. What is the value of Ksp for BaCrO4? A. 2.1 ´ 10–10 B. 8.6 ´ 10–10 C. 1.4 ´ 10–5 D. 1.5 ´ 10–5 E. 2.9 ´ 10–5
36. What is the concentration of Ag+ in a saturated solution of Ag2CrO4 if Ksp = 1 ´ 10-12? A. 1 ´ 10-12 M B. 5 ´ 10-5 M C. 6 ´ 10-5 M D. 1 ´ 10-4 M E. 2 ´ 10-4 M
37. What is the concentration of Hg22+ in a saturated solution of Hg2Cl2 if Ksp = 1 ´ 10-18? A. 1 ´ 10-9 M B. 2 ´ 10-9 M C. 3 ´ 10-7 M D. 6 ´ 10-7 M E. 1 ´ 10-6 M
38. What is the water solubility of AgCl (Ksp = 1.8 ´ 10-10) in 0.25 M NaCl? A. 4.5 ´ 10-11 M B. 7.2 ´ 10-10 M C. 1.8 ´ 10-9 M D. 1.7 ´ 10-7 M E. 1.3 ´ 10-5 M
39. What is the water solubility of BaCO3 (Ksp = 2.6 ´ 10-9) in a solution containing 0.20 M CO32-(aq)? A. 5.2 ´ 10-10 M B. 1.3 ´ 10-8 M C. 5.1 ´ 10-5 M D. 2.5 ´ 10-4 M E. 1.0 ´ 10-5 M
40. What is the water solubility of PbI2 (Ksp = 8.4 ´ 10-9) in 0.15 M KI? A. 3.7 ´ 10-7 M B. 5.6 ´ 10-7 M C. 1.9 ´ 10-6 M D. 1.1 ´ 10-4 M E. 6.1 ´ 10-4 M
41. What is the water solubility of BaF2(s) in 0.033 M KF(aq) at 25C? The Ksp of BaF2 is 1.8 ´ 10–7 at 25C. A. 2.7 ´ 10–6 M B. 5.5 ´ 10–6 M C. 4.1 ´ 10–5 M D. 1.1 ´ 10–4 M E. 1.7 ´ 10–4 M
42. What is the molar solubility of Fe(OH)3(s) in a solution that is buffered at pH 2.50 at 25C? The Ksp of Fe(OH)3 is 6.3 ´ 10–38 at 25C. A. 6.9 ´ 10–28 M B. 2.0 ´ 10–26 M C. 1.3 ´ 10–13 M D. 2.0 ´ 10–3 M E. 5.0 ´ 102 M
43. An aqueous solution contains 0.010 M Br– and 0.010 M I–. If Ag+ is added until AgBr(s) just begins to precipitate, what are the concentrations of Ag+ and I–? (Ksp of AgBr = 5.4 ´ 10–13, Ksp of AgI = 8.5 ´ 10–17) A. [Ag+] = 5.4 ´ 10–11 M, [I–] = 1.0 ´ 10–2 M B. [Ag+] = 8.5 ´ 10–15 M, [I–] = 1.0 ´ 10–2 M C. [Ag+] = 5.4 ´ 10–11 M, [I–] = 1.6 ´ 10–6 M D. [Ag+] = 8.5 ´ 10–15 M, [I–] = 6.4 ´ 101 M E. [Ag+] = 8.5 ´ 10–15 M, [I–] = 1.6 ´ 10–6 M
44. A solution contains 0.10 M Ca2+ and 0.10 M Mg2+. The pH of the solution is raised without changing the volume of the solution. What percentage of Mg2+ remains in solution when Ca(OH)2(s) first begins to precipitate? (Ksp of Ca(OH)2 = 4.0 ´ 10-6 and Ksp of Mg(OH)2 = 7.1 ´ 10-12) A. 1.8 ´ 10-4 % B. 4.0 ´ 10-4 % C. 6.3 ´ 10-3 % D. 7.2 ´ 10-3 % E. 1.4 ´ 10-2 %
45. What is the net ionic equation for the reaction of iron(III) hydroxide with a strong acid? A. 3OH-(aq) + 3H+(aq) ® 3H2O(l) B. Fe3+(aq) + 3OH-(aq) + 3H+(aq) ® Fe3+(aq) + 3H2O(l) C. Fe(OH)3(s) + 3H+(aq) ® Fe3+(aq) + 3H2O(l) D. Fe3+(aq) + 3H+(aq) ® Fe3+(aq) + 3H2O(l) E. Fe(OH)3(s) + 3H+(aq) ® FeH3(s) + 3OH-(aq)
46. What is the net ionic equation for the reaction of Ca3(PO4)2 with a strong acid? A. Ca3(PO4)2(s) + 6H+(aq) ® 3CaH3(s) + 2PO43-(aq) B. Ca3(PO4)2(s) + 6H+(aq) ® 3Ca2+(aq) + 2H3PO4(aq) C. 2PO43-(aq) + 2H+(aq) ® H2PO4-(aq) D. 3Ca2+(aq) + 2PO43-(aq) + 6H+(aq) ® 2H3PO4(aq) + 3Ca2+(aq) E. Ca3(PO4)2(s) + 16H+(aq) ® Ca3P2(s) + 8H2O(l)
47. Determine the equilibrium constant, K, for the following reaction, Mg(OH)2(s) + 2H+(aq) ® Mg2+(aq) + 2H2O(l) given the Ksp of Mg(OH)2 is 6 ´ 10-12 and Kw is 1.0 ´ 10-14. A. 6 ´ 10-26 B. 6 ´ 10-12 C. 1 ´ 10-14 D. 6 ´ 1016 E. 2 ´ 10-3
48. Determine the equilibrium constant, K, for the following reaction, Ca(OH)2(s) + 2H+(aq) ® Ca2+(aq) + 2H2O(l) given the Ksp of Ca(OH)2 is 4.0 ´ 10-6 and Kw is 1.0 ´ 10--14. A. 4.0 ´ 10-20 B. 4.0 ´ 10-6 C. 1.0 ´ 10-14 D. 2.5 ´ 10-9 E. 4.0 ´ 1022
49. What is the net ionic equation for the reaction of KCl with a strong base? A. KCl(s) + OH-(aq) ® KOH(aq) + Cl-(aq) B. KCl(aq) + OH-(aq) ® KOH(aq) + Cl-(aq) C. K+(aq) + OH-(aq) ® KOH(s) D. Cl-(aq) + OH-(aq) ® HOCl(aq) E. No reaction will occur.
50. What is the net ionic equation for the reaction of ZnS with a strong acid? A. ZnS(s) + 2H+(aq) ® Zn2+(aq) + H2S(aq) B. Zn2+(aq) + 2H+(aq) ® ZnH2(aq) C. S2-(aq) + 2H+(aq) ® H2S(aq) D. S2-(aq) + 2H+(aq) ® H2S(s) E. No reaction will occur.
51. What is the net ionic equation for the reaction of an aqueous solution of AgNO3 and NH3? A. AgNO3(s) + NH3(aq) ® AgNH3(s) + NO3-(aq) B. Ag+(aq) + NH3(aq) ® AgNH3(s) C. Ag+(aq) + 2NH3(aq) ® Ag(NH3)2+(aq) D. AgNO3(s) + 2NH3(aq) ® Ag(NH3)2+(aq) + NO3-(aq) E. No reaction will occur.
52. Given the following reactions, AgBr(s) Ag+(aq) + Br–(aq) + Ag (aq) + 2 CN–(aq) Ag(CN)2–(aq)
Ksp = 5.4 ´ 10–13 Kf = 1.2 ´ 1021
determine the equilibrium constant for the reaction below. AgBr(s) + 2 CN–(aq) Ag(CN)2–(aq) + Br–(aq)
A. 4.5 ´ 10–34 B. 1.5 ´ 10–9 C. 6.5 ´ 108 D. 1.2 ´ 1021 E. 2.2 ´ 1033 53. Consider the reaction Cu(OH)2(s) + 4NH3(aq)
Cu(NH3)42+(aq) + 2OH-(aq)
K = 4 ´ 10-7
If the Kf for Cu(NH3)42+ is 1 ´ 1012, what is the value of Ksp for Cu(OH)2?
A. 4 ´ 10-19 B. 2 ´ 10-13 C. 5 ´ 10-12 D. 4 ´ 10-7 E. 4 ´ 105 54. A solution containing an unknown metal ion is analyzed by qualitative analysis. Addition of chloride has no effect on the solution. Addition of H2S at pH 0.5 results in a precipitate. What group of cations is present? A. Group I B. Group II C. Group III D. Group IV E. Group V
55. Which of the following metals will precipitate as chloride salts: Ag+, Pb2+, Ca2+, K+, and Cu2+? A. Ag+ B. Pb2+, Ca2+, and Cu2+ C. Ag+, K+, and Cu2+ D. Ag+ and Pb2+ E. Ca2+ and Cu2+
56. In the qualitative analysis scheme, Mg2+ is a group IV cation. What anion is used to precipitate Mg2+? A. OHB. ClC. PO43D. S2E. CO32-
57. Which of the boxes below represents solid MX2 in equilibrium with M2+ (squares) and X- (circles)?
A.
B.
C.
D.
E.
58. A soluble ionic species, SX, is dissolved in water. An excess of slightly soluble solid, MX, is added to the solution. Which of the boxes below represents MX in equilibrium with M+ (squares) and X- (circles) in the presence of SX?
A.
B.
C.
D.
E.
59. A solution containing an unknown metal ion is analyzed by qualitative analysis. Addition of chloride has no effect on the solution. Addition of H2S at pH 0.5 results in no precipitate. Addition of H2S at pH 9.0 results in no precipitate. What group(s) of cations may be present? A. Groups I&II B. Groups II&V C. Group III D. Group IV E. Groups IV&V
60. The addition of an anion to dissolve one precipitate from a mixture of precipitates is known as _____ . A. solubility product B. common ion effect C. complex formation D. solution formation E. selective precipitation
Chapter 15--Complex Ion and Precipitation Equilibria Key 1. Which of the following equations represents the formation constant of Ti(NH3)63+? A. Ti3+(aq) + (NH3)6(aq) Ti(NH3)63+(aq) B. Ti(NH3)63+(aq) Ti3+(aq) + 6NH3(aq) 3+ C. Ti (aq) + NH3(aq) Ti(NH3)3+(aq) D. 6Ti3+(aq) + 6NH3(aq) Ti(NH3)63+(aq) 3+ E. Ti (aq) + 6NH3(aq) Ti(NH3)63+(aq)
2. The formation constant for Zn(NH3)42+ is 3.6 ´ 108. What is the ratio of Zn(NH3)42+ to Zn2+ in 0.010 M NH3? A. 3.6 B. 1.4 ´ 103 C. 3.6 ´ 106 D. 1.4 ´ 103 E. 3.6 ´ 1010
3. The formation constant for Ag(CN)2-(aq) is 2 ´ 1020. At what cyanide ion concentration is [Ag+] = [ Ag(CN)2-]? A. 3 ´ 10-41 M B. 5 ´ 10-21 M C. 7 ´ 10-11 M D. 6 ´ 10-6 M E. 1 ´ 1010 M
4. Ksp stands for A. saturated precipitate. B. soluble precipitate. C. soluble particle. D. saturated precipitate. E. solubility product.
5. What is the chemical equation that describes the dissolution of calcium iodate in water? A. CaI2(s) Ca2+(aq) + 2I-(aq) 2+ B. Ca (aq) + IO3-(aq) CaIO3(s) C. Ca(IO3)2(s) Ca(s) + I2(aq) + 3O2(g) D. Ca(IO3)2(s) Ca2+(aq) + 2IO3-(aq) E. Ca(IO3)2(s) CaI2(aq) + 3O2(g)
6. What is the chemical equation that describes the dissolution of nickel(II) hydroxide in water? A. Ni(OH)2(s) Ni2+(aq) + 2OH-(aq) 2+ B. Ni (aq) + 3OH (aq) Ni(OH)2(s) + C. Ni2OH(s) + H (aq) 2Ni+(aq) + H2O(l) + D. Ni(OH)2(s) + 2H (aq) Ni2+(aq) + 2H2O(l) E. Ni(OH)2(aq) + H2O(l) Ni(OH)3(s) + H+(aq)
7. What is the chemical equation that describes the dissolution of silver oxalate, Ag2(C2O4), in water? A. Ag2(C2O4)(s) Ag22+(aq) + C2O42-(aq) B. Ag2(C2O4)(s) 2Ag+(aq) + C2O42-(aq) C. Ag2(C2O4)(s) Ag2+(aq) + C2O4-(aq) + 2D. 2Ag (aq) + C2O4 (aq) Ag2(C2O4)(s) E. Ag2+(aq) + C2O4-(aq) Ag2(C2O4)(s)
8. Which of the following equations is the solubility product of Cu(II) hydroxide, Cu(OH)2?
A.
B. C. D. E.
9. What is the equilibrium constant expression for the dissolution of solid lead oxalate, PbC2O4, in water? A. B. Ksp = [Pb]2 [C2O4]2
C.
D. E.
10. What is the equilibrium constant expression for the dissolution of solid calcium fluoride in water? A. Ksp = [Ca2+][F-] B. Ksp = [Ca2+]2[F-] C. Ksp = [Ca2+][F-]2
D.
E.
11. Consider the equilibrium of PbCl2(s) in water. PbCl2(s) Pb2+(aq) + 2Cl-(aq) What is the effect of adding NaCl(aq) to the equilibrium solution? A. The sodium ion reduces the Pb2+ to Pb(s). B. PbCl2 solubility increases due to the common-ion effect. C. PbNa2(s) precipitates. D. The NaCl(aq) has no effect on the system. E. PbCl2(s) precipitates until equilibrium is reestablished.
12. Consider the equilibrium of Ca(OH)2(s) in water. Ca(OH)2(s) Ca2+(aq) + 2OH-(aq) What is the effect of raising the pH of the solution? A. Ca2+(aq) is reduced to Ca(s). B. The concentration of hydronium ion increases. C. The concentration of Ca2+ increases as Ca(OH)2 dissolves. D. Ca(OH)2(s) precipitates until equilibrium is reestablished. E. Hydroxide ion is reduced to H2(g).
13. Consider the equilibrium of lead sulfide, PbS, in water. PbS(s) Pb2+(aq) + S2-(aq) What is the effect of adding PbS(s) to the solution? A. The PbS(s) has no effect on the system. B. More Pb2+ and S2- will form. C. More Pb2+ will form, but the concentration of S2- will remain constant. D. More S2- will form, but the concentration of Pb2+ will remain constant. E. The concentrations of Pb2+ and S2- will decrease.
14. Consider the equilibrium of magnesium fluoride, MgF2(s), in water. MgF2(s) Mg2+(aq) + 2F-(aq) What is the effect of adding perchloric acid (i.e. decreasing the pH of solution)? A. The acid will have no effect on the magnesium fluoride equilibrium. B. The F- will be protonated, resulting in more MgF2(s) dissolving. C. More MgF2(s) will be formed, thus reducing Mg2+ and F- concentrations in solution. D. Magnesium perchlorate, Mg(ClO4)2(s), will precipitate. E. Magnesium hydride, MgH2(s), will precipitate.
15. The Ksp of barium chromate is 1.2 ´ 10-10. What is the concentration of Ba2+ in equilibrium with BaCrO4(s) if [CrO42-] = 4.3 ´ 10-3 M? A. 5.1 ´ 10-13 M B. 2.8 ´ 10-8 M C. 1.1 ´ 10-5 M D. 1.7 ´ 10-4 M E. 2.5 ´ 10-3 M
16. The Ksp of calcium iodate is 7.1 ´ 10-7. What is the concentration of Ca2+ in equilibrium with Ca(IO3)2(s) if [IO3-] = 3.3 ´ 10-2 M? A. 3.3 ´ 10-7 M B. 2.2 ´ 10-5 M C. 8.1 ´ 10-5 M D. 6.5 ´ 10-4 M E. 7.1 ´ 10-3 M
17. The Ksp of aluminum hydroxide, Al(OH)3, is 2 ´ 10-31. What pH is required to limit the Al3+ concentration to less than or equal to 1 ´ 10-10 M? A. 3.6 B. 6.4 C. 7.1 D. 7.8 E. 11.5
18. The Ksp of calcium phosphate is 1 ´ 10-33. What is the concentration of Ca2+ in equilibrium with Ca3(PO4)2(s) if [PO43-] = 1 ´ 10-5 M? A. 1 ´ 10-28 M B. 1 ´ 10-23 M C. 1 ´ 10-11 M D. 5 ´ 10-10 M E. 2 ´ 10-8 M
19. The Ksp of Fe(OH)3(s) is 3 ´ 10-39. What concentration of Fe3+ can exist in solution at pH 3.0? A. 3 ´ 10-4 M B. 3 ´ 10-6 M C. 3 ´ 10-21 M D. 3 ´ 10-24 M E. 3 ´ 10-30 M
20. The Ksp of BaSO4 is 1.1 ´ 10–10 at 25C. What mass of BaSO4 (molar mass = 233.4 g/mol) will dissolve in 1.0 L of water at 25C? A. 2.6 ´ 10-8 g B. 4.5 ´ 10-8 g C. 1.0 ´ 10-5 g D. 1.6 ´ 10-4 g E. 2.4 ´ 10-3 g
21. For AgCl, Ksp = 1.8 ´ 10-10. What will occur if 250 mL of 1.5 ´ 10-3 M NaCl is mixed with 250 mL of 2.0 ´ 10-7 M AgNO3? A. Q > Ksp. A precipitate will form. B. Ksp > Q. A precipitate will form. C. Q = Ksp. No precipitate will form. D. Q > Ksp. No precipitate will form. E. Ksp > Q. No precipitate will form.
22. For PbCl2, Ksp = 1.7 ´ 10-5. What will occur if 250 mL of 0.12 M Pb(NO3)2 is mixed with 250 mL of 0.070 M NaCl? A. Q > Ksp. A precipitate will form. B. Ksp > Q. A precipitate will form. C. Q = Ksp. No precipitate will form. D. Q > Ksp. No precipitate will form. E. Ksp > Q. No precipitate will form.
23. For Ca(OH)2, Ksp = 4.0 ´ 10-6. What will occur if 1.0 L of 0.100 M Ca(NO3)2 is prepared in a solution that is buffered at pH 12.50? A. Q > Ksp. A precipitate will form. B. Ksp > Q. A precipitate will form. C. Q = Ksp. No precipitate will form. D. Q > Ksp. No precipitate will form. E. Ksp > Q. No precipitate will form.
24. For Mg(OH)2, Ksp = 6 ´ 10-12. What will occur if 1.0 L of 0.010 M Mg(NO3)2 is prepared at pH 10.00? A. Q > Ksp. A precipitate will form. B. Ksp > Q. A precipitate will form. C. Q = Ksp. No precipitate will form. D. Q > Ksp. No precipitate will form. E. Ksp > Q. No precipitate will form.
25. What is the water solubility of AgI (Ksp = 1 ´ 10-16, MM = 234.8 g/mol) in moles per liter? A. 1 ´ 10-32 M B. 1 ´ 10-16 M C. 1 ´ 10-8 M D. 2 ´ 10-6 M E. 4 ´ 10-11 M
26. What is the water solubility of PbI2 (Ksp = 8.4 ´ 10-9, MM = 461 g/mol) in moles per liter? A. 8.4 ´ 10-9 M B. 1.9 ´ 10-7 M C. 9.1 ´ 10-5 M D. 1.3 ´ 10-3 M E. 2.0 ´ 10-3 M
27. What is the water solubility of Hg2Br2 (Ksp = 6 ´ 10-23) in moles per liter? A. 8 ´ 10-12 M B. 1 ´ 10-8 M C. 2 ´ 10-8 M D. 4 ´ 10-8 M E. 3 ´ 10-6 M
28. What is the water solubility of AgCl (Ksp = 1.8 ´ 10-10, MM = 143.4 g/mol) in grams per liter? A. 1.3 ´ 10-12 g/L B. 2.6 ´ 10-8 g/L C. 9.4 ´ 10-8 g/L D. 9.6 ´ 10-4 g/L E. 1.9 ´ 10-3 g/L
29. The Ksp of BaSO4 is 1.1 ´ 10–10 at 25C. What mass of BaSO4 (molar mass = 233.4 g/mol) will dissolve in 1.0 L of water at 25C? A. 2.6 ´ 10-8 g B. 4.5 ´ 10-8 g C. 1.0 ´ 10-5 g D. 1.6 ´ 10-4 g E. 2.4 ´ 10-3 g
30. What volume of water is needed to completely dissolve 100 g of AgBr (Ksp = 5 ´ 10-13, 187.8 g/mol)? A. 1 ´ 10-2 L B. 8 ´ 105 L C. 1 ´ 108 L D. 3 ´ 1010 L E. 1 ´ 1012 L
31. What volume of water is needed to completely dissolve 1.0 g of Ag2CrO4 (Ksp = 8.0 ´ 10-12, MM = 331.8 g/mol)? A. 0.94 L B. 1.1 L C. 1.7 L D. 24 L E. 2.6 ´ 105 L
32. The Ksp of Ca(OH)2 is 5.5 ´ 10–5 at 25C. What is the concentration of OH–(aq) in a saturated solution of Ca(OH)2(aq)? A. 1.9 ´ 10–3 M B. 7.4 ´ 10–3 M C. 2.4 ´ 10–2 M D. 4.0 ´ 10–2 M E. 4.8 ´ 10–2 M
33. A saturated solution of lead(II) sulfate can be prepared by diluting 0.0101 g of PbSO4 to 250 mL. What is the Ksp of lead(II) sulfate? A. 8.3 ´ 10-12 B. 1.6 ´ 10-9 C. 1.8 ´ 10-8 D. 4.5 ´ 10-8 E. 8.4 ´ 10-7
34. At pH 10.0, only 0.019 g of MgCl2 will dissolve per 1 L of solution. What is the Ksp of magnesium hydroxide, Mg(OH)2? A. 2 ´ 10-13 B. 2 ´ 10-12 C. 2 ´ 10-10 D. 2 ´ 10-8 E. 2 ´ 10-6
35. The solubility of BaCrO4(s) in water is 3.7 milligrams in 1.0 L at 25C. What is the value of Ksp for BaCrO4? A. 2.1 ´ 10–10 B. 8.6 ´ 10–10 C. 1.4 ´ 10–5 D. 1.5 ´ 10–5 E. 2.9 ´ 10–5
36. What is the concentration of Ag+ in a saturated solution of Ag2CrO4 if Ksp = 1 ´ 10-12? A. 1 ´ 10-12 M B. 5 ´ 10-5 M C. 6 ´ 10-5 M D. 1 ´ 10-4 M E. 2 ´ 10-4 M
37. What is the concentration of Hg22+ in a saturated solution of Hg2Cl2 if Ksp = 1 ´ 10-18? A. 1 ´ 10-9 M B. 2 ´ 10-9 M C. 3 ´ 10-7 M D. 6 ´ 10-7 M E. 1 ´ 10-6 M
38. What is the water solubility of AgCl (Ksp = 1.8 ´ 10-10) in 0.25 M NaCl? A. 4.5 ´ 10-11 M B. 7.2 ´ 10-10 M C. 1.8 ´ 10-9 M D. 1.7 ´ 10-7 M E. 1.3 ´ 10-5 M
39. What is the water solubility of BaCO3 (Ksp = 2.6 ´ 10-9) in a solution containing 0.20 M CO32-(aq)? A. 5.2 ´ 10-10 M B. 1.3 ´ 10-8 M C. 5.1 ´ 10-5 M D. 2.5 ´ 10-4 M E. 1.0 ´ 10-5 M
40. What is the water solubility of PbI2 (Ksp = 8.4 ´ 10-9) in 0.15 M KI? A. 3.7 ´ 10-7 M B. 5.6 ´ 10-7 M C. 1.9 ´ 10-6 M D. 1.1 ´ 10-4 M E. 6.1 ´ 10-4 M
41. What is the water solubility of BaF2(s) in 0.033 M KF(aq) at 25C? The Ksp of BaF2 is 1.8 ´ 10–7 at 25C. A. 2.7 ´ 10–6 M B. 5.5 ´ 10–6 M C. 4.1 ´ 10–5 M D. 1.1 ´ 10–4 M E. 1.7 ´ 10–4 M
42. What is the molar solubility of Fe(OH)3(s) in a solution that is buffered at pH 2.50 at 25C? The Ksp of Fe(OH)3 is 6.3 ´ 10–38 at 25C. A. 6.9 ´ 10–28 M B. 2.0 ´ 10–26 M C. 1.3 ´ 10–13 M D. 2.0 ´ 10–3 M E. 5.0 ´ 102 M
43. An aqueous solution contains 0.010 M Br– and 0.010 M I–. If Ag+ is added until AgBr(s) just begins to precipitate, what are the concentrations of Ag+ and I–? (Ksp of AgBr = 5.4 ´ 10–13, Ksp of AgI = 8.5 ´ 10–17) A. [Ag+] = 5.4 ´ 10–11 M, [I–] = 1.0 ´ 10–2 M B. [Ag+] = 8.5 ´ 10–15 M, [I–] = 1.0 ´ 10–2 M C. [Ag+] = 5.4 ´ 10–11 M, [I–] = 1.6 ´ 10–6 M D. [Ag+] = 8.5 ´ 10–15 M, [I–] = 6.4 ´ 101 M E. [Ag+] = 8.5 ´ 10–15 M, [I–] = 1.6 ´ 10–6 M
44. A solution contains 0.10 M Ca2+ and 0.10 M Mg2+. The pH of the solution is raised without changing the volume of the solution. What percentage of Mg2+ remains in solution when Ca(OH)2(s) first begins to precipitate? (Ksp of Ca(OH)2 = 4.0 ´ 10-6 and Ksp of Mg(OH)2 = 7.1 ´ 10-12) A. 1.8 ´ 10-4 % B. 4.0 ´ 10-4 % C. 6.3 ´ 10-3 % D. 7.2 ´ 10-3 % E. 1.4 ´ 10-2 %
45. What is the net ionic equation for the reaction of iron(III) hydroxide with a strong acid? A. 3OH-(aq) + 3H+(aq) ® 3H2O(l) B. Fe3+(aq) + 3OH-(aq) + 3H+(aq) ® Fe3+(aq) + 3H2O(l) C. Fe(OH)3(s) + 3H+(aq) ® Fe3+(aq) + 3H2O(l) D. Fe3+(aq) + 3H+(aq) ® Fe3+(aq) + 3H2O(l) E. Fe(OH)3(s) + 3H+(aq) ® FeH3(s) + 3OH-(aq)
46. What is the net ionic equation for the reaction of Ca3(PO4)2 with a strong acid? A. Ca3(PO4)2(s) + 6H+(aq) ® 3CaH3(s) + 2PO43-(aq) B. Ca3(PO4)2(s) + 6H+(aq) ® 3Ca2+(aq) + 2H3PO4(aq) C. 2PO43-(aq) + 2H+(aq) ® H2PO4-(aq) D. 3Ca2+(aq) + 2PO43-(aq) + 6H+(aq) ® 2H3PO4(aq) + 3Ca2+(aq) E. Ca3(PO4)2(s) + 16H+(aq) ® Ca3P2(s) + 8H2O(l)
47. Determine the equilibrium constant, K, for the following reaction, Mg(OH)2(s) + 2H+(aq) ® Mg2+(aq) + 2H2O(l) given the Ksp of Mg(OH)2 is 6 ´ 10-12 and Kw is 1.0 ´ 10-14. A. 6 ´ 10-26 B. 6 ´ 10-12 C. 1 ´ 10-14 D. 6 ´ 1016 E. 2 ´ 10-3
48. Determine the equilibrium constant, K, for the following reaction, Ca(OH)2(s) + 2H+(aq) ® Ca2+(aq) + 2H2O(l) given the Ksp of Ca(OH)2 is 4.0 ´ 10-6 and Kw is 1.0 ´ 10--14. A. 4.0 ´ 10-20 B. 4.0 ´ 10-6 C. 1.0 ´ 10-14 D. 2.5 ´ 10-9 E. 4.0 ´ 1022
49. What is the net ionic equation for the reaction of KCl with a strong base? A. KCl(s) + OH-(aq) ® KOH(aq) + Cl-(aq) B. KCl(aq) + OH-(aq) ® KOH(aq) + Cl-(aq) C. K+(aq) + OH-(aq) ® KOH(s) D. Cl-(aq) + OH-(aq) ® HOCl(aq) E. No reaction will occur.
50. What is the net ionic equation for the reaction of ZnS with a strong acid? A. ZnS(s) + 2H+(aq) ® Zn2+(aq) + H2S(aq) B. Zn2+(aq) + 2H+(aq) ® ZnH2(aq) C. S2-(aq) + 2H+(aq) ® H2S(aq) D. S2-(aq) + 2H+(aq) ® H2S(s) E. No reaction will occur.
51. What is the net ionic equation for the reaction of an aqueous solution of AgNO3 and NH3? A. AgNO3(s) + NH3(aq) ® AgNH3(s) + NO3-(aq) B. Ag+(aq) + NH3(aq) ® AgNH3(s) C. Ag+(aq) + 2NH3(aq) ® Ag(NH3)2+(aq) D. AgNO3(s) + 2NH3(aq) ® Ag(NH3)2+(aq) + NO3-(aq) E. No reaction will occur.
52. Given the following reactions, AgBr(s) Ag+(aq) + Br–(aq) + Ag (aq) + 2 CN–(aq) Ag(CN)2–(aq)
Ksp = 5.4 ´ 10–13 Kf = 1.2 ´ 1021
determine the equilibrium constant for the reaction below. AgBr(s) + 2 CN–(aq) Ag(CN)2–(aq) + Br–(aq)
A. 4.5 ´ 10–34 B. 1.5 ´ 10–9 C. 6.5 ´ 108 D. 1.2 ´ 1021 E. 2.2 ´ 1033 53. Consider the reaction Cu(OH)2(s) + 4NH3(aq)
Cu(NH3)42+(aq) + 2OH-(aq)
K = 4 ´ 10-7
If the Kf for Cu(NH3)42+ is 1 ´ 1012, what is the value of Ksp for Cu(OH)2?
A. 4 ´ 10-19 B. 2 ´ 10-13 C. 5 ´ 10-12 D. 4 ´ 10-7 E. 4 ´ 105 54. A solution containing an unknown metal ion is analyzed by qualitative analysis. Addition of chloride has no effect on the solution. Addition of H2S at pH 0.5 results in a precipitate. What group of cations is present? A. Group I B. Group II C. Group III D. Group IV E. Group V
55. Which of the following metals will precipitate as chloride salts: Ag+, Pb2+, Ca2+, K+, and Cu2+? A. Ag+ B. Pb2+, Ca2+, and Cu2+ C. Ag+, K+, and Cu2+ D. Ag+ and Pb2+ E. Ca2+ and Cu2+
56. In the qualitative analysis scheme, Mg2+ is a group IV cation. What anion is used to precipitate Mg2+? A. OHB. ClC. PO43D. S2E. CO32-
57. Which of the boxes below represents solid MX2 in equilibrium with M2+ (squares) and X- (circles)?
A.
B.
C.
D.
E.
58. A soluble ionic species, SX, is dissolved in water. An excess of slightly soluble solid, MX, is added to the solution. Which of the boxes below represents MX in equilibrium with M+ (squares) and X- (circles) in the presence of SX?
A.
B.
C.
D.
E.
59. A solution containing an unknown metal ion is analyzed by qualitative analysis. Addition of chloride has no effect on the solution. Addition of H2S at pH 0.5 results in no precipitate. Addition of H2S at pH 9.0 results in no precipitate. What group(s) of cations may be present? A. Groups I&II B. Groups II&V C. Group III D. Group IV E. Groups IV&V
60. The addition of an anion to dissolve one precipitate from a mixture of precipitates is known as _____ . A. solubility product B. common ion effect C. complex formation D. solution formation E. selective precipitation
Chapter 16--Spontaneity of Reaction 1. The following processes occur spontaneously at 25C. Which of these processes is/are endothermic? 1. 2. 3.
NH4NO3 dissolving in water (which is accompanied by a cooling of the water). the expansion of a real gas into a vacuum (which is accompanied by a cooling of the gas). liquid water in an ice cube tray freezing into ice after being placed in a freezer.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 2. The third law of thermodynamics states that A. there is no disorder in a perfect crystal at 0 K. B. in a spontaneous process, the entropy of the universe increases. C. the total entropy of the universe is always increasing. D. the total mass of the universe is constant. E. mass and energy are conserved in all chemical reactions.
3. Which of the following statements concerning entropy is/are CORRECT? 1. 2. 3.
The entropy of a substance increases when converted from a liquid to a solid. The entropy of a substance decreases as its temperature increases. All substances have positive standard molar entropies at temperatures above 0 K.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 4. The second law of thermodynamics states that A. in a spontaneous process, there is a net increase in entropy, taking into account both the system and the surroundings. B. there is no disorder in a perfect crystal at 0 K. C. the total energy of the universe is always increasing. D. the total energy of the universe is constant. E. mass and energy are conserved in all chemical reactions.
5. Predict the signs of DS and DH for the evaporation of water at 295 K. A. DH = 0 and DS > 0 B. DH > 0 and DS < 0 C. DH > 0 and DS > 0 D. DH < 0 and DS > 0 E. DH < 0 and DS < 0
6. Predict the sign of DS and DH for the combustion of gasoline at 500 K. A. DH > 0 and DS > 0 B. DH > 0 and DS = 0 C. DH < 0 and DS > 0 D. DH < 0 and DS < 0 E. DH < 0 and DS = 0
7. When a real gas is compressed from low pressure to a higher pressure, its temperature increases. Predict the signs of DH and DS. A. DH < 0 and DS < 0 B. DH < 0 and DS > 0 C. DH > 0 and DS < 0 D. DH > 0 and DS > 0 E. DH < 0 and DS = 0
8. All of the following statements are false for the freezing of water at 273 K EXCEPT A. DH < 0. B. DH > 0. C. DH = 0. D. DS = 0. E. DS > 0.
9. The dissolution of ammonium nitrate occurs spontaneously in water at 25C. As NH4NO3 dissolves, the temperature of the water decreases. What are the signs of DH, DS, and DG for this process? A. DH > 0, DS < 0, DG > 0 B. DH > 0, DS > 0, DG > 0 C. DH > 0, DS > 0, DG < 0 D. DH < 0, DS < 0, DG < 0 E. DH < 0, DS > 0, DG > 0
10. Diluting concentrated sulfuric acid with water can be dangerous. The temperature of the solution can increase rapidly. What are the signs of DH, DS, and DG for this process? A. DH < 0, DS > 0, DG < 0 B. DH < 0, DS < 0, DG < 0 C. DH < 0, DS > 0, DG > 0 D. DH > 0, DS > 0, DG < 0 E. DH > 0, DS < 0, DG > 0
11. If a chemical reaction has a negative change in entropy, DS, then A. the reaction is endothermic. B. the reaction is spontaneous. C. the equilibrium constant is greater than 1. D. there is an increase in the order of the system. E. the change in Gibbs free energy, DG, is negative.
12. All of the following events result in an increase in entropy EXCEPT A. the melting of candle wax. B. the combustion of carbon. C. the dissolution of sodium chloride in water. D. the evaporation of ethanol. E. the formation of N2O4(g) from NO2(g).
13. All of the following statements are true EXCEPT A. DS > 0 for systems that become more disorderly. B. reactions are spontaneous when DG < 0. C. reactions are spontaneous when DH < 0. D. DH < 0 for exothermic reactions. E. generally, nature tends to move from more ordered to more random states.
14. Which of the following linear chain alcohols is likely to have the largest standard molar entropy in the liquid state? A. CH3OH B. CH3CH2OH C. CH3CH2CH2OH D. CH3CH2CH2CH2OH E. CH3CH2CH2CH2CH2OH
15. Calculate DS for the following reaction, H2(g) + Br2(l) ® 2HBr(g) given S[H2(g)] = +131 J/molK, S[Br2(l)] = +152 J/molK, and S[HBr(g)] = +199 J/molK. A. -84 J/K B. +84 J/K C. +115 J/K D. +482 J/K E. +681 J/K
16. Calculate DS for the evaporation of water, given S[H2O(l)] = +69.9 J/molK and S[H2O(g)] = +188.7 J/molK. A. -258.6 J/K B. -118.8 J/K C. 0.0000 J/K D. +118.8 J/K E. +258.6 J/K
17. Calculate DS for the following reaction, 2SO2(g) + O2(g) ® 2SO3(g) given S[SO2(g)] = 248.2 J/molK, S[O2(g)] = 205.1 J/molK, and S[SO3(g)] = 256.8 J/molK. A. –196.5 J/K B. –94.0 J/K C. –187.9 J/K D. +187.9 J/K E. +196.5 J/K
18. Which of the following reactions is likely to have the most positive change in entropy? A. N2(g) + 3H2(g) ® 2NH3(g) B. N2(g) + 2O2(g) ® 2NO2(g) C. 2C(s) + O2(g) ® 2CO(g) D. C(s) + O2(g) ® CO2(g) E. CaO(s) + CO2(g) ® CaCO3(s)
19. For which of the following reactions will the entropy of the system decrease? A. 2NH3(g) ® N2(g) + 3H2(g) B. 2C(s) + O2(g) ® 2CO(g) C. CaCO3(s) ® CaO(s) + CO2(g) D. 2NO2(g) ® N2O4(g) E. NaOH(s) ® Na+(aq) + OH–(aq)
20. All of the following statements are true EXCEPT A. for a given material, the gas has greater entropy than the solid. B. the entropy for gaseous elements is zero at 298 K. C. at 0 K, an ordered pure crystalline solid has an entropy of zero. D. for a given material, the liquid has greater entropy than the solid. E. increasing the temperature of a substance increases its entropy.
21. All of the following statements are true EXCEPT A. a reaction is spontaneous if DG < 0. B. if DG = 0, the system is at equilibrium. C. if DG = 0, then DS = DH. D. if DG > 0, then a reaction is not spontaneous. E. DG is referred to as Gibbs free energy.
22. If a chemical reaction is exothermic, but not spontaneous, which of the following must be true? A. DG > 0, DS > 0 and DH > 0 B. DG < 0, DS > 0 and DH > 0 C. DG > 0, DS < 0 and DH > 0 D. DG < 0, DS < 0 and DH < 0 E. DG > 0, DS < 0 and DH < 0
23. If DG is positive at all temperatures, then which of the following statements must be true? A. DH < 0 and DS < 0 B. DH < 0 and DS = 0 C. DH < 0 and DS > 0 D. DH > 0 and DS > 0 E. DH > 0 and DS < 0
24. If DG is negative at all temperatures, then which of the following statements must be true? A. DH < 0 and DS < 0 B. DH < 0 and DS = 0 C. DH < 0 and DS > 0 D. DH > 0 and DS > 0 E. DH > 0 and DS < 0
25. When ammonium nitrate dissolves spontaneously in water the temperature of the solution decreases. Which statement is true for this system? A. DH < 0 and DS < 0 B. DH < 0 and DS = 0 C. DH = 0 and DS > 0 D. DH > 0 and DS > 0 E. DH > 0 and DS < 0
26. At what temperatures will a reaction be spontaneous if DH = +62.4 kJ and DS = +301 J/K? A. All temperatures below 207 K. B. All temperatures above 207 K. C. Temperatures between 179 K and 235 K. D. The reaction will be spontaneous at any temperature. E. The reaction will never be spontaneous.
27. Which of the following are conditions for the standard Gibbs free energy? 1. 2. 3.
Ions or molecules in solution are present at one molar concentrations. Gases are present at one atmosphere partial pressure. The temperature of the system is 273 K.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 28. What is the correct form for the Gibbs-Helmholtz equation? A. DG = DH - TDS B. DG = DH + TDS C. DH = DG - TDS D. DS = DH - TDG E. DG = DS - TDH
29. Calculate DG for the reaction below at 25.0C CS2(g) + 3Cl2(g) ® S2Cl2(g) + CCl4(g) given DH = –231.1 kJ and DS = –287.6 J/K. A. –518.7 kJ B. –316.9 kJ C. –145.3 kJ D. –56.5 kJ E. +56.5 kJ
30. What is DG at 555C for a reaction with DH = -152 kJ and DS = 79.1 J/K? A. -4370 kJ B. -217 kJ C. -196 kJ D. +196 kJ E. + 643 kJ
31. Ammonia is synthesized from nitrogen and hydrogen gases at a temperature of 475C. N2(g) + 3H2(g) ® 2NH3(g) If DH = -92.2 kJ and DS = -0.1987 kJ/K, what is DG for the reaction at 575C? A. -260.7 kJ B. -186.6 kJ C. -92.0 kJ D. +2.2 kJ E. +76.3 kJ
32. Ammonia is synthesized from nitrogen and hydrogen gases. N2(g) + 3H2(g) ® 2NH3(g) If DH = -92.2 kJ and DS = -0.1987 kJ/K, at what temperature will DG = 0? A. 0.00216 K B. 18.3 K C. 92.0 K D. 464 K E. 672 K
33. Nitric oxide can be made from the reaction of oxygen and nitrogen gases. O2(g) + N2(g) ® 2NO(g) If DG = 165.5 kJ, and DH = 180.4 kJ, what is DS at 325C? A. 0.0125 kJ/K B. 0.0249 kJ/K C. 0.0458 kJ/K D. 0.142 kJ/K E. 1.02 kJ/K
34. Calculate DG for the reaction below at 25.0C. P4O10(s) + 6H2O(l) ® 4H3PO4(l)
Species P4O10(s) H2O(l) H3PO4(l)
S (J/molK ) –2984.0 –285.8 –1279.0
228.9 69.95 110.5
A. –6119.1 kJ B. –632.8 kJ C. –355.6 kJ D. +210.6 kJ E. +6119.1 kJ 35. Calculate DG for the following reaction at 298 K. S(s) + O2(g) ® SO2(g)
Substance S(s) O2(g) SO2(g)
S (kJ/molK) 0.0 0.0 -296.8
+0.0318 +0.2050 +0.2481
A. -300.2 kJ B. -296.8 kJ C. -85.1 kJ D. +29.3 kJ E. +145.7 kJ 36. At 298 K, all of the following substances have a standard free energy of formation of zero EXCEPT ____. A. Br2(g) B. I2(s) C. S8(s) D. Cl2(g) E. Hg(l)
37. Calculate DS for the dissolution of BaCl2 in water at 25C. BaCl2(s) ® Ba2+(aq) + 2Cl-(aq)
Substance BaCl2(s) Ba2+(aq) Cl-(aq)
-810.4 -560.8 -131.2
-858.6 -537.6 -167.2
A. -0.20 kJ/K B. -0.14 kJ/K C. -0.0020 kJ/K D. +0.12 kJ/K E. +0.29 kJ/K 38. Calculate DS for the dissociation of dinitrogen tetraoxide at 25C. N2O4(g) ® 2NO2(g)
Substance NO2(g) N2O4(g)
+51.3 +97.9
+33.2 +9.2
A. -2.10 kJ/K B. -0.550 kJ/K C. -0.208 kJ/K D. +0.176 kJ/K E. +2.10 kJ/K
39. Calculate
at 325 K for ethylene, C2H4(g), given the thermodynamic data below.
Substance C(s) H2(g) C2H4(g)
A. -47.1 kJ B. -19.0 kJ C. +39.1 kJ D. +52.1 kJ E. +69.6 kJ
S (kJ/molK) 0.0 0.0 +52.3
+0.0057 +0.1306 +0.2195
40. For the decomposition of hydrogen peroxide to water and oxygen, H2O2(g) ® H2O(g) + O2(g) DH = -106 kJ, and DS = +0.0580 kJ/K. In what temperature range is the reaction spontaneous (DG < 0)? A. The temperature must be greater than 1.83 ´ 103 K. B. The temperature must be less than 1.83 ´ 103 K. C. The temperature must be between 225 K and 1.83 ´ 103 K. D. DG is always less than zero. E. DG is never less than zero.
41. In what temperature range is DG greater than zero for the formation of NH4Cl(s) from NH3(g) and HCl(g)? NH3(g) + HCl(g) ® NH4Cl(s) For the reaction, DH = -176.0 kJ and DS = -0.2845 kJ/K. A. The temperature must be greater than 619 K. B. The temperature must be less than 619 K. C. The temperature must be exactly 619 K. D. DG is always greater than zero. E. DG is never greater than zero.
42. In a gas phase reaction, what is the effect of increasing reactant or product pressure on the standard Gibbs free energy, DG? A. DG increases due to decreased entropy. B. DG decreases due to decreased entropy. C. DG increases due to increased enthalpy. D. DG may either increase or decrease. E. DG is unchanged.
43. The standard free energy change associated with the dissolution of ammonium nitrate in water is –6.73 kJ at 298 K. NH4NO3(s) NH4NO3(aq) What is the equilibrium constant for the reaction? (R = 8.31 ´ 10-3 kJ/K) A. 1.9 ´ 10–3 B. 6.6 ´ 10–2 C. 1.0 D. 15 E. 5.2 ´ 102
44. For the dimerization of nitrogen dioxide, DG = -4.7 kJ at 25C. 2NO2(g) ® N2O4(g) Calculate DG for this reaction if the partial pressures of NO2 and N2O4 are both 0.50 atm. (R = 8.31 ´ 10-3 kJ/K) A. -6.4 kJ B. -4.8 kJ C. -4.7 kJ D. -4.6 kJ E. -3.0 kJ
45. The standard free energy change for a chemical reaction is +13.3 kJ. What is the equilibrium constant for the reaction at 125C? (R = 8.31 ´ 10-3 kJ/K) A. 2.8 ´ 10–6 B. 2.0 ´ 10–5 C. 4.7 ´ 10–3 D. 1.8 ´ 10–2 E. 2.1 ´ 102
46. At 25C, the acid dissociation constant for formic acid is 6.9 ´ 10-4. Calculate DG for this reaction. (R = 8.31 ´ 10-3 kJ/K) A. -18.0 kJ B. -1.51 kJ C. +1.51 kJ D. +18.0 kJ E. +41.4 kJ
47. At 25C, the equilibrium constant for the following reaction, Cu2+(aq) + Zn(s) ® Cu(s) + Zn2+(aq) is 1.9 ´ 1037. Calculate DG for this reaction. (R = 8.31 ´ 10-3 kJ/K) A. -213 kJ B. -145 kJ C. -57.3 kJ D. +57.3 kJ E. +213 kJ
48. If DG = 0, then A. K = 0. B. K < 0. C. K < 1. D. K = 1. E. K < -1.
49. For the decomposition of calcium carbonate, DG = +130.4 kJ at 25C. CaCO3(s) ® CaO(s) + CO2(g) Calculate the partial pressure of CO2 if DG = 0.0 kJ. (R = 8.31 ´ 10-3 kJ/K) A. 1.3 ´ 10-23 atm B. 2.7 ´ 102 atm C. 8.2 ´ 10-5 atm D. 5.3 ´ 101 atm E. 7.4 ´ 1022 atm
50. Given that, at 25C, C(s) + O2(g) ® CO2(g) C(s) +
DG = -394.4 kJ DG = -137.2 kJ
O2(g) ® CO(g)
calculate DG for the reaction below. CO(g) +
O2(g) ® CO2(g)
A. -531.6 kJ B. -257.2 kJ C. -171.5 kJ D. +141.2 kJ E. +531.6 kJ 51. Given that S(g) + O2(g) ® SO2(g) 2S(g) + 3O2(g) ® 2SO3(g)
calculate of the following reaction: SO2(g) + 1/2O2(g) ® SO3(g)
A. –1042.2 kJ B. –71.0 kJ C. +2.47 kJ D. +71.0 kJ E. +1042.2 kJ 52. For a chemical reaction, DG and DG are equal A. when products and reactants are in standard state concentrations. B. when K = 0. C. when K > 1. D. for a system at equilibrium. E. when the entropy change is zero.
DG = –300.1 kJ DG = –742.1 kJ
53. J. Willard Gibbs, professor at Yale from 1871 to 1903, A. was a researcher, experimenting with high energy reactions. B. was a theoretician, who studied chemical thermodynamics. C. received no salary for the first ten years. D. both a and c. E. both b and c.
54. Rubber elasticity is an entropic phenomenon. As a rubber band is stretched the entropy of the system A. is reduced. B. is increased. C. remains unchanged. D. breaks the polymer chain. E. causes chemical bonds to form.
Chapter 16--Spontaneity of Reaction Key
1. The following processes occur spontaneously at 25C. Which of these processes is/are endothermic? 1. 2. 3.
NH4NO3 dissolving in water (which is accompanied by a cooling of the water). the expansion of a real gas into a vacuum (which is accompanied by a cooling of the gas). liquid water in an ice cube tray freezing into ice after being placed in a freezer.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 2. The third law of thermodynamics states that A. there is no disorder in a perfect crystal at 0 K. B. in a spontaneous process, the entropy of the universe increases. C. the total entropy of the universe is always increasing. D. the total mass of the universe is constant. E. mass and energy are conserved in all chemical reactions.
3. Which of the following statements concerning entropy is/are CORRECT? 1. 2. 3.
The entropy of a substance increases when converted from a liquid to a solid. The entropy of a substance decreases as its temperature increases. All substances have positive standard molar entropies at temperatures above 0 K.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 4. The second law of thermodynamics states that A. in a spontaneous process, there is a net increase in entropy, taking into account both the system and the surroundings. B. there is no disorder in a perfect crystal at 0 K. C. the total energy of the universe is always increasing. D. the total energy of the universe is constant. E. mass and energy are conserved in all chemical reactions.
5. Predict the signs of DS and DH for the evaporation of water at 295 K. A. DH = 0 and DS > 0 B. DH > 0 and DS < 0 C. DH > 0 and DS > 0 D. DH < 0 and DS > 0 E. DH < 0 and DS < 0
6. Predict the sign of DS and DH for the combustion of gasoline at 500 K. A. DH > 0 and DS > 0 B. DH > 0 and DS = 0 C. DH < 0 and DS > 0 D. DH < 0 and DS < 0 E. DH < 0 and DS = 0
7. When a real gas is compressed from low pressure to a higher pressure, its temperature increases. Predict the signs of DH and DS. A. DH < 0 and DS < 0 B. DH < 0 and DS > 0 C. DH > 0 and DS < 0 D. DH > 0 and DS > 0 E. DH < 0 and DS = 0
8. All of the following statements are false for the freezing of water at 273 K EXCEPT A. DH < 0. B. DH > 0. C. DH = 0. D. DS = 0. E. DS > 0.
9. The dissolution of ammonium nitrate occurs spontaneously in water at 25C. As NH4NO3 dissolves, the temperature of the water decreases. What are the signs of DH, DS, and DG for this process? A. DH > 0, DS < 0, DG > 0 B. DH > 0, DS > 0, DG > 0 C. DH > 0, DS > 0, DG < 0 D. DH < 0, DS < 0, DG < 0 E. DH < 0, DS > 0, DG > 0
10. Diluting concentrated sulfuric acid with water can be dangerous. The temperature of the solution can increase rapidly. What are the signs of DH, DS, and DG for this process? A. DH < 0, DS > 0, DG < 0 B. DH < 0, DS < 0, DG < 0 C. DH < 0, DS > 0, DG > 0 D. DH > 0, DS > 0, DG < 0 E. DH > 0, DS < 0, DG > 0
11. If a chemical reaction has a negative change in entropy, DS, then A. the reaction is endothermic. B. the reaction is spontaneous. C. the equilibrium constant is greater than 1. D. there is an increase in the order of the system. E. the change in Gibbs free energy, DG, is negative.
12. All of the following events result in an increase in entropy EXCEPT A. the melting of candle wax. B. the combustion of carbon. C. the dissolution of sodium chloride in water. D. the evaporation of ethanol. E. the formation of N2O4(g) from NO2(g).
13. All of the following statements are true EXCEPT A. DS > 0 for systems that become more disorderly. B. reactions are spontaneous when DG < 0. C. reactions are spontaneous when DH < 0. D. DH < 0 for exothermic reactions. E. generally, nature tends to move from more ordered to more random states.
14. Which of the following linear chain alcohols is likely to have the largest standard molar entropy in the liquid state? A. CH3OH B. CH3CH2OH C. CH3CH2CH2OH D. CH3CH2CH2CH2OH E. CH3CH2CH2CH2CH2OH
15. Calculate DS for the following reaction, H2(g) + Br2(l) ® 2HBr(g) given S[H2(g)] = +131 J/molK, S[Br2(l)] = +152 J/molK, and S[HBr(g)] = +199 J/molK. A. -84 J/K B. +84 J/K C. +115 J/K D. +482 J/K E. +681 J/K
16. Calculate DS for the evaporation of water, given S[H2O(l)] = +69.9 J/molK and S[H2O(g)] = +188.7 J/molK. A. -258.6 J/K B. -118.8 J/K C. 0.0000 J/K D. +118.8 J/K E. +258.6 J/K
17. Calculate DS for the following reaction, 2SO2(g) + O2(g) ® 2SO3(g) given S[SO2(g)] = 248.2 J/molK, S[O2(g)] = 205.1 J/molK, and S[SO3(g)] = 256.8 J/molK. A. –196.5 J/K B. –94.0 J/K C. –187.9 J/K D. +187.9 J/K E. +196.5 J/K
18. Which of the following reactions is likely to have the most positive change in entropy? A. N2(g) + 3H2(g) ® 2NH3(g) B. N2(g) + 2O2(g) ® 2NO2(g) C. 2C(s) + O2(g) ® 2CO(g) D. C(s) + O2(g) ® CO2(g) E. CaO(s) + CO2(g) ® CaCO3(s)
19. For which of the following reactions will the entropy of the system decrease? A. 2NH3(g) ® N2(g) + 3H2(g) B. 2C(s) + O2(g) ® 2CO(g) C. CaCO3(s) ® CaO(s) + CO2(g) D. 2NO2(g) ® N2O4(g) E. NaOH(s) ® Na+(aq) + OH–(aq)
20. All of the following statements are true EXCEPT A. for a given material, the gas has greater entropy than the solid. B. the entropy for gaseous elements is zero at 298 K. C. at 0 K, an ordered pure crystalline solid has an entropy of zero. D. for a given material, the liquid has greater entropy than the solid. E. increasing the temperature of a substance increases its entropy.
21. All of the following statements are true EXCEPT A. a reaction is spontaneous if DG < 0. B. if DG = 0, the system is at equilibrium. C. if DG = 0, then DS = DH. D. if DG > 0, then a reaction is not spontaneous. E. DG is referred to as Gibbs free energy.
22. If a chemical reaction is exothermic, but not spontaneous, which of the following must be true? A. DG > 0, DS > 0 and DH > 0 B. DG < 0, DS > 0 and DH > 0 C. DG > 0, DS < 0 and DH > 0 D. DG < 0, DS < 0 and DH < 0 E. DG > 0, DS < 0 and DH < 0
23. If DG is positive at all temperatures, then which of the following statements must be true? A. DH < 0 and DS < 0 B. DH < 0 and DS = 0 C. DH < 0 and DS > 0 D. DH > 0 and DS > 0 E. DH > 0 and DS < 0
24. If DG is negative at all temperatures, then which of the following statements must be true? A. DH < 0 and DS < 0 B. DH < 0 and DS = 0 C. DH < 0 and DS > 0 D. DH > 0 and DS > 0 E. DH > 0 and DS < 0
25. When ammonium nitrate dissolves spontaneously in water the temperature of the solution decreases. Which statement is true for this system? A. DH < 0 and DS < 0 B. DH < 0 and DS = 0 C. DH = 0 and DS > 0 D. DH > 0 and DS > 0 E. DH > 0 and DS < 0
26. At what temperatures will a reaction be spontaneous if DH = +62.4 kJ and DS = +301 J/K? A. All temperatures below 207 K. B. All temperatures above 207 K. C. Temperatures between 179 K and 235 K. D. The reaction will be spontaneous at any temperature. E. The reaction will never be spontaneous.
27. Which of the following are conditions for the standard Gibbs free energy? 1. 2. 3.
Ions or molecules in solution are present at one molar concentrations. Gases are present at one atmosphere partial pressure. The temperature of the system is 273 K.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 28. What is the correct form for the Gibbs-Helmholtz equation? A. DG = DH - TDS B. DG = DH + TDS C. DH = DG - TDS D. DS = DH - TDG E. DG = DS - TDH
29. Calculate DG for the reaction below at 25.0C CS2(g) + 3Cl2(g) ® S2Cl2(g) + CCl4(g) given DH = –231.1 kJ and DS = –287.6 J/K. A. –518.7 kJ B. –316.9 kJ C. –145.3 kJ D. –56.5 kJ E. +56.5 kJ
30. What is DG at 555C for a reaction with DH = -152 kJ and DS = 79.1 J/K? A. -4370 kJ B. -217 kJ C. -196 kJ D. +196 kJ E. + 643 kJ
31. Ammonia is synthesized from nitrogen and hydrogen gases at a temperature of 475C. N2(g) + 3H2(g) ® 2NH3(g) If DH = -92.2 kJ and DS = -0.1987 kJ/K, what is DG for the reaction at 575C? A. -260.7 kJ B. -186.6 kJ C. -92.0 kJ D. +2.2 kJ E. +76.3 kJ
32. Ammonia is synthesized from nitrogen and hydrogen gases. N2(g) + 3H2(g) ® 2NH3(g) If DH = -92.2 kJ and DS = -0.1987 kJ/K, at what temperature will DG = 0? A. 0.00216 K B. 18.3 K C. 92.0 K D. 464 K E. 672 K
33. Nitric oxide can be made from the reaction of oxygen and nitrogen gases. O2(g) + N2(g) ® 2NO(g) If DG = 165.5 kJ, and DH = 180.4 kJ, what is DS at 325C? A. 0.0125 kJ/K B. 0.0249 kJ/K C. 0.0458 kJ/K D. 0.142 kJ/K E. 1.02 kJ/K
34. Calculate DG for the reaction below at 25.0C. P4O10(s) + 6H2O(l) ® 4H3PO4(l)
Species P4O10(s) H2O(l) H3PO4(l)
S (J/molK ) –2984.0 –285.8 –1279.0
228.9 69.95 110.5
A. –6119.1 kJ B. –632.8 kJ C. –355.6 kJ D. +210.6 kJ E. +6119.1 kJ 35. Calculate DG for the following reaction at 298 K. S(s) + O2(g) ® SO2(g)
Substance S(s) O2(g) SO2(g)
S (kJ/molK) 0.0 0.0 -296.8
+0.0318 +0.2050 +0.2481
A. -300.2 kJ B. -296.8 kJ C. -85.1 kJ D. +29.3 kJ E. +145.7 kJ 36. At 298 K, all of the following substances have a standard free energy of formation of zero EXCEPT ____. A. Br2(g) B. I2(s) C. S8(s) D. Cl2(g) E. Hg(l)
37. Calculate DS for the dissolution of BaCl2 in water at 25C. BaCl2(s) ® Ba2+(aq) + 2Cl-(aq)
Substance BaCl2(s) Ba2+(aq) Cl-(aq)
-810.4 -560.8 -131.2
-858.6 -537.6 -167.2
A. -0.20 kJ/K B. -0.14 kJ/K C. -0.0020 kJ/K D. +0.12 kJ/K E. +0.29 kJ/K 38. Calculate DS for the dissociation of dinitrogen tetraoxide at 25C. N2O4(g) ® 2NO2(g)
Substance NO2(g) N2O4(g)
+51.3 +97.9
+33.2 +9.2
A. -2.10 kJ/K B. -0.550 kJ/K C. -0.208 kJ/K D. +0.176 kJ/K E. +2.10 kJ/K
39. Calculate
at 325 K for ethylene, C2H4(g), given the thermodynamic data below.
Substance C(s) H2(g) C2H4(g)
A. -47.1 kJ B. -19.0 kJ C. +39.1 kJ D. +52.1 kJ E. +69.6 kJ
S (kJ/molK) 0.0 0.0 +52.3
+0.0057 +0.1306 +0.2195
40. For the decomposition of hydrogen peroxide to water and oxygen, H2O2(g) ® H2O(g) + O2(g) DH = -106 kJ, and DS = +0.0580 kJ/K. In what temperature range is the reaction spontaneous (DG < 0)? A. The temperature must be greater than 1.83 ´ 103 K. B. The temperature must be less than 1.83 ´ 103 K. C. The temperature must be between 225 K and 1.83 ´ 103 K. D. DG is always less than zero. E. DG is never less than zero.
41. In what temperature range is DG greater than zero for the formation of NH4Cl(s) from NH3(g) and HCl(g)? NH3(g) + HCl(g) ® NH4Cl(s) For the reaction, DH = -176.0 kJ and DS = -0.2845 kJ/K. A. The temperature must be greater than 619 K. B. The temperature must be less than 619 K. C. The temperature must be exactly 619 K. D. DG is always greater than zero. E. DG is never greater than zero.
42. In a gas phase reaction, what is the effect of increasing reactant or product pressure on the standard Gibbs free energy, DG? A. DG increases due to decreased entropy. B. DG decreases due to decreased entropy. C. DG increases due to increased enthalpy. D. DG may either increase or decrease. E. DG is unchanged.
43. The standard free energy change associated with the dissolution of ammonium nitrate in water is –6.73 kJ at 298 K. NH4NO3(s) NH4NO3(aq) What is the equilibrium constant for the reaction? (R = 8.31 ´ 10-3 kJ/K) A. 1.9 ´ 10–3 B. 6.6 ´ 10–2 C. 1.0 D. 15 E. 5.2 ´ 102
44. For the dimerization of nitrogen dioxide, DG = -4.7 kJ at 25C. 2NO2(g) ® N2O4(g) Calculate DG for this reaction if the partial pressures of NO2 and N2O4 are both 0.50 atm. (R = 8.31 ´ 10-3 kJ/K) A. -6.4 kJ B. -4.8 kJ C. -4.7 kJ D. -4.6 kJ E. -3.0 kJ
45. The standard free energy change for a chemical reaction is +13.3 kJ. What is the equilibrium constant for the reaction at 125C? (R = 8.31 ´ 10-3 kJ/K) A. 2.8 ´ 10–6 B. 2.0 ´ 10–5 C. 4.7 ´ 10–3 D. 1.8 ´ 10–2 E. 2.1 ´ 102
46. At 25C, the acid dissociation constant for formic acid is 6.9 ´ 10-4. Calculate DG for this reaction. (R = 8.31 ´ 10-3 kJ/K) A. -18.0 kJ B. -1.51 kJ C. +1.51 kJ D. +18.0 kJ E. +41.4 kJ
47. At 25C, the equilibrium constant for the following reaction, Cu2+(aq) + Zn(s) ® Cu(s) + Zn2+(aq) is 1.9 ´ 1037. Calculate DG for this reaction. (R = 8.31 ´ 10-3 kJ/K) A. -213 kJ B. -145 kJ C. -57.3 kJ D. +57.3 kJ E. +213 kJ
48. If DG = 0, then A. K = 0. B. K < 0. C. K < 1. D. K = 1. E. K < -1.
49. For the decomposition of calcium carbonate, DG = +130.4 kJ at 25C. CaCO3(s) ® CaO(s) + CO2(g) Calculate the partial pressure of CO2 if DG = 0.0 kJ. (R = 8.31 ´ 10-3 kJ/K) A. 1.3 ´ 10-23 atm B. 2.7 ´ 102 atm C. 8.2 ´ 10-5 atm D. 5.3 ´ 101 atm E. 7.4 ´ 1022 atm
50. Given that, at 25C, C(s) + O2(g) ® CO2(g) C(s) +
DG = -394.4 kJ DG = -137.2 kJ
O2(g) ® CO(g)
calculate DG for the reaction below. CO(g) +
O2(g) ® CO2(g)
A. -531.6 kJ B. -257.2 kJ C. -171.5 kJ D. +141.2 kJ E. +531.6 kJ 51. Given that S(g) + O2(g) ® SO2(g) 2S(g) + 3O2(g) ® 2SO3(g)
calculate of the following reaction: SO2(g) + 1/2O2(g) ® SO3(g)
A. –1042.2 kJ B. –71.0 kJ C. +2.47 kJ D. +71.0 kJ E. +1042.2 kJ 52. For a chemical reaction, DG and DG are equal A. when products and reactants are in standard state concentrations. B. when K = 0. C. when K > 1. D. for a system at equilibrium. E. when the entropy change is zero.
DG = –300.1 kJ DG = –742.1 kJ
53. J. Willard Gibbs, professor at Yale from 1871 to 1903, A. was a researcher, experimenting with high energy reactions. B. was a theoretician, who studied chemical thermodynamics. C. received no salary for the first ten years. D. both a and c. E. both b and c.
54. Rubber elasticity is an entropic phenomenon. As a rubber band is stretched the entropy of the system A. is reduced. B. is increased. C. remains unchanged. D. breaks the polymer chain. E. causes chemical bonds to form.
Chapter 17--Electrochemistry 1. If a reactant loses electrons it is A. reduced. B. oxidized. C. disproportionated. D. the cathode. E. a catalyst.
2. The anode is where A. electrolyte concentrations are lowest. B. electrolyte concentrations are highest. C. the Nernst potential is measured. D. anions are produced. E. oxidation occurs.
3. All of the following statements concerning voltaic cells are true EXCEPT A. a salt bridge allows cations and anions to move between the half-cells. B. electrons flow from the cathode to the anode in the external circuit. C. oxidation occurs at the anode. D. a voltaic cell can be used as a source of energy. E. a voltaic cell consists of two-half cells.
4. The following reaction occurs spontaneously. 2 Fe(s) + 3 Cl2(aq) ® 2 Fe3+(aq) + 6 Cl–(aq) Write the balanced reduction half-reaction. A. Fe(s) ® Fe3+(aq) + e– B. Fe(s) + 3 e– ® Fe3+(aq) C. Fe(s) + 3/2 Cl2(aq) ® FeCl3(aq) D. Cl2(aq) ® 2 Cl–(aq) + 2 e– E. Cl2(aq) + 2 e– ® 2 Cl–(aq)
5. The following reaction occurs spontaneously. 2Al(s) + 6H+(aq) ® 2Al3+(aq) + 3H2(g) Write the balanced oxidation half-reaction. A. Al(s) ® Al3+(aq) + 3eB. Al3+(aq) + 3e- ® Al(s) C. H2(g) ® 2H+(aq) + 2eD. 2H+(aq) + 2e- ® H2(g) E. 2H+(aq) ® H2(g) + 2e-
6. Given the following half-reactions, Pb(s) + 2I-(aq) ® PbI2(s) + 2eI2(aq) + 2e- ® 2I-(aq) write the balanced overall reaction. A. Pb(s) + I2(aq) + 2e- ® PbI (s) + I-(aq) B. Pb(s) + I2(aq) ® PbI2(s) C. Pb(s) + Cl2(aq) ® Pb2+(aq) + 2I-(aq) D. 2Pb(s) + 4e- ® 2PbI(s) + I2(aq) E. Pb(s) + e- + 2I-(aq) ® PbI2(aq)
7. Given the following half-reactions, 2IO3-(aq) + 12H+(aq) + 10e- ® I2(s) + 6H2O(l) Fe2+(aq) ® Fe3+(aq) + ewrite the balanced overall reaction. A. 2IO3-(aq) + Fe2+(aq) ® I2(s) + Fe3+(aq) B. 2IO3-(aq) + 10Fe2+(aq) ® I2(s) + 10Fe3+(aq) C. 2IO3-(aq) + 12H+(aq) + Fe2+(aq) ® FeI2(s) + 6H2O(l) D. 2IO3-(aq) + 12H+(aq) + Fe2+(aq) ® I2(s) + Fe3+(aq) + 6H2O(l) E. 2IO3-(aq) + 12H+(aq) + 10Fe2+(aq) ® I2(s) + 10Fe3+(aq) + 6H2O(l)
8. Write a balanced net ionic equation for the overall reaction represented by the cell notation below. Ag | AgI | I– || Fe2+, Fe3+ | Pt A. Fe3+(aq) + Ag(s) + I–(aq) ® Fe2+(aq) + AgI(s) B. Fe2+(aq) + Ag(s) + I–(aq) ® Fe3+(aq) + AgI(s) C. Fe3+(aq) + AgI(s) ® Fe2+(aq) + Ag(s) + I–(aq) D. Fe2+(aq) + AgI(s) ® Fe3+(aq) + Ag(s) + I–(aq) E. Fe3+(aq) + 2I–(aq) ® Fe2+(aq) + I2(s)
9. Write a balanced chemical equation for the overall reaction represented below. Pt | Fe2+, Fe3+ || Cl- | AgCl | Ag A. AgCl(s) + Fe3+(aq) + 2e- ® Ag(s) + Fe2+(aq) + Cl-(aq) B. AgCl(s) + Fe3+(aq) ® Ag(s) + Fe2+(aq) + Cl-(aq) C. AgCl(s) + Fe2+(aq) ® Ag(s) + Fe3+(aq) + Cl-(aq) D. Ag(s) + Fe3+(aq) + Cl-(aq) ® AgCl(s) + Fe2+(aq) E. Ag(s) + Fe2+(aq) + Cl-(aq) ® AgCl(s) + Fe3+(aq)
10. Write a balanced chemical equation for the overall reaction represented below. Cu | Cu2+ || Cl- | Hg2Cl2 | Hg A. 2Hg(l) + Cu2+(aq) + 2Cl-(aq) ® Hg2Cl2(s) + Cu(s) B. Hg2Cl2(s) + Cu(s) ® 2Hg(l) + Cu2+(aq) + 2Cl-(aq) C. Hg2Cl2(s) ® 2Hg(l) + Cl2(aq) D. Hg2Cl2(s) + Cu(s) ® 2Hg(l) + Cu2+(aq) + Cl2(aq) E. 2Hg(l) + Cl2(aq) ® Hg2Cl2(s)
11. What is the correct cell notation for the reaction below? Cd2+(aq) + Ni(s) ® Cd(s) + Ni2+(aq) A. Cd2+ | Cd || Ni | Ni2+ B. Cd | Cd2+ || Ni2+ | Ni C. Ni | Ni2+ || Cd2+ | Cd D. Ni | Cd2+ || Ni2+ | Cd E. Cd2+ | Cd || Ni | Ni2+
12. What is a correct cell notation for a voltaic cell based on the reaction below? Cu2+(aq) + Pb(s) + SO42–(aq) ® Cu(s) + PbSO4(s) A. Pb | PbSO4 || Cu2+ || Cu B. Cu | Cu2+ || SO42– | PbSO4 | Pb C. Cu | Cu2+, SO42– | PbSO4 | Pb D. Cu | Cu2+, SO42– || PbSO4 || Pb E. Pb | PbSO4 | SO42– || Cu2+ || Cu
13. Which species is the best oxidizing agent? A. bromine B. sodium ion C. potassium metal D. fluoride ion E. carbon dioxide
14. Which species is the best reducing agent? A. sodium ion B. sulfide ion C. chlorine D. sodium metal E. silver metal
15. Use the standard reduction potentials below to determine which element or ion is the best reducing agent. NO3–(aq) + 4H+(aq) + 3e– ® NO(g) + 2H2O(l) Pd2+(aq) + 2e– ® Pd(s) 2H+(aq) + 2e– ® H2(g)
E = +0.955 V E = +0.90 V E = 0.00 V
A. Pd2+(aq) B. Pd(s) C. H+(aq) D. H2(g) E. NO3–(aq)
16. The standard reduction potential, oxidation potential. What is the value of
, of Zn2+(aq) is -0.762 V. Write the equation for the standard ?
A. Zn2+(aq) + 2e- ® Zn(s)
= -0.762 V
B. Zn2+(aq) ® Zn(s) + 2e-
= +0.762 V
C. Zn2+(aq) ® Zn(s) + 2e-
= -0.762 V
D. Zn(s) + 2e- ® Zn2+(aq)
= -0.762 V
E. Zn(s) ® Zn2+(aq) + 2e-
= +0.762 V
17. The standard oxidation potential, , of Cu(s) is -0.339 V. The oxidation product is the copper(II) ion. Write the equation for the standard reduction potential. What is the value of the standard reduction potential, ? A. Cu2+(aq) ® Cu(s) + 2e-
= -0.339 V
B. Cu2+(aq) + 2e- ® Cu(s)
= +0.339 V
C. Cu(s) + 2e- ® Cu2+(aq)
= -0.339 V
D. Cu(s) ® Cu2+(aq) + 2e-
= +0.339 V
E. Cu(s) + 2e- ® Cu2+(aq)
= +0.339 V
18. Which of the following species are likely to behave as oxidizing agents: Li(s), H2(g), MnO4–(aq), and Cl–(aq)? A. Li(s) only B. MnO4–(aq) only C. H2(g) and Cl–(aq) D. Li(s) and MnO4–(aq) E. Cl–(aq) only
19. Which of the following species are reducing agents: K, H+, MnO4-, Cl2, Sn4+? A. K only B. K and Cl2 C. H+ and Sn4+ D. MnO4- and Cl2 E. Cl2 and Sn4+
20. Write the overall chemical reaction and calculate E for the following reduction and oxidation half-reactions. 2H+(aq) + 2e- ® H2(g) = 0.000 V Zn(s) ® Zn2+(aq) + 2e= +0.762 V
A. Zn2+(aq) + H2(g) ® Zn(s) + 2H+(aq) B. Zn2+(aq) + H2(g) ® Zn(s) + 2H+(aq) C. Zn(s) + 2H+(aq) ® Zn2+(aq) + H2(g) D. Zn(s) + 2H+(aq) ® Zn2+(aq) + H2(g) E. Zn(s) + 2H+(aq) ® Zn2+(aq) + H2(g)
E = +0.381 V E = -0.762 V E = +0.762 V E = -0.762 V E = -0.381 V
21. Given the following two half-reactions, write the overall balanced reaction in the direction in which it is spontaneous and calculate the standard cell potential. Ga3+(aq) + 3e– ® Ga(s) = –0.53 V Sn4+(aq) + 2e– ® Sn2+(aq) = +0.15 V
A. 2Ga3+(aq) + 3Sn2+(aq) ® 2Ga(s) + 3Sn4+(aq) B. 3Ga3+(aq) + 2Sn2+(aq) ® 3Ga(s) + 2Sn4+(aq) C. 2Ga(s) + 3Sn4+(aq) ® 2Ga3+(aq) + 3Sn2+(aq) D. 3Ga(s) + 2Sn4+(aq) ® 3Ga3+(aq) + 2Sn2+(aq) E. 2Ga(s) + 3Sn4+(aq) ® 2Ga3+(aq) + 3Sn2+(aq)
E = +0.68 V E = –1.89 V E = +0.68 V E = +0.38 V E = +1.89 V
22. Using the reduction half-reactions, write an overall reaction in the direction that is spontaneous. Calculate E. AgI(s) + e- ® Ag(s) + I-(aq) = -0.152 V Co3+(aq) + e- ® Co2+(aq) = +1.953 V
A. Co2+(aq) + AgI(s) ® Co3+(aq) + Ag(s) + I-(aq) B. Co2+(aq) + Ag(s) + I-(aq) ® Co3+(aq) + AgI(s) C. Co3+(aq) + Ag(s) + I-(aq) ® Co2+(aq) + AgI(s) D. Co3+(aq) + Ag(s) + I-(aq) ® Co2+(aq) + AgI(s) E. Co3+(aq) + AgI(s) ® Co2+(aq) + Ag(s) + I-(aq)
E = -2.105 V E = -1.801 V E = +1.801 V E = +2.105 V E = -2.105 V
23. Using the reduction half-reactions, write an overall reaction in the direction that it is spontaneous. Calculate E. MnO4-(aq) + 8H+(aq) + 5e- ® 2Mn2+(aq) + 4H2O(l) = +1.512 V Fe3+(aq) + e- ® Fe2+(aq) = +0.769 V
A. 5Fe3+(aq) + Mn2+(aq) + 4H2O(l) ® 5Fe2+(aq) + MnO4-(aq) + 8H+(aq) B. 5Fe2+(aq) + Mn2+(aq) + 4H2O(l) ® 5Fe3+(aq) + MnO4-(aq) + 8H+(aq) C. MnO4-(aq) + 5Fe2+(aq) ® Mn2+(aq) + 5Fe3+(aq) E = -2.281 V D. MnO4-(aq) + 8H+(aq) + 5Fe2+(aq) ® Mn2+(aq) + 5Fe3+(aq) + 4H2O(l) E. MnO4-(aq) + 8H+(aq) + 5Fe3+(aq) ® Mn2+(aq) + 5Fe2+(aq) + 4H2O(l)
E = +2.333 V E = -0.743 V E = +0.743 V E = +2.281 V
24. The cell voltage, E, is ____ for a reaction taking place in a voltaic cell. A. reductive B. oxidative C. negative D. zero E. positive
25. Which of the following half-reactions has been assigned a standard reduction potential of 0.00 V? A. 2H+(aq) + 2e- ® H2(g) B. Ag+(aq) + e- ® Ag(s) C. Hg2Cl2(s) + 2e- ® 2Hg(l) + 2Cl-(aq) D. Au3+(aq) + 3e- ® Au(s) E. Cu2+(aq) + 2e- ® Cu(s)
26. Calculate E for an electrochemical cell based on the following overall reaction: 2H+(aq) + Cd(s) ® H2(g) + Cd2+(aq) if (H+) = 0.000 V and A. -0.402 V B. -0.201 V C. 0.000 V D. +0.201 V E. +0.402 V
(Cd2+) = -0.402 V.
27. Calculate E for the electrochemical cell below, Ag | AgCl| Cl– || Cu2+ | Cu given the following standard reduction potentials. Cu2+(aq) + 2e– ® Cu(s) AgCl(s) + e– ® Ag(s) + Cl–(aq)
E = +0.337 V E = +0.222 V
A. –0.115 V B. –0.107 V C. +0.115 V D. +0.452 V E. +0.559 V 28. Use the standard reduction potentials below to determine which of the following metals can be oxidized by 1.0 M HCl: Al, Zn, Sn, Cu. Al3+(aq) + 3e- ® Al(s) = -1.68 V Zn2+(aq) + 2e- ® Zn(s) = -0.762 V Sn2+(aq) + 2e- ® Sn(s) = -0.141 V Cu2+(aq) + 2e- ® Cu(s) = +0.339 V
A. Al only B. Cu only C. Al and Cu D. Zn and Sn E. Al, Zn, and Sn 29. Calculate E for the electrochemical cell below, Pb |PbCl2 | Cl– || Fe3+, Fe2+ | Pt given the following reduction half-reactions. Pb2+(aq) + 2e– ® Pb(s) PbCl2(s) + 2e– ® Pb(s) + 2Cl–(aq) Fe3+(aq) + e– ® Fe2+(aq) Fe2+(aq) + e– ® Fe(s)
A. –0.504 V B. –0.062 V C. +0.504 V D. +1.038 V E. +1.604 V
E = –0.126 V E = –0.267 V E = +0.771 V E = –0.44 V
30. The standard reduction potentials for F2(g) and Br2(l) are +2.889 V and 1.077 V, respectively. Write the chemical equation and calculate E for a voltaic cell based on fluorine and bromine. A. F2(g) + 2Br-(aq) ® 2F-(aq) + Br2(g) E = +1.812 V B. F2(g) + Br2(g) ® 2F-(aq) + 2Br-(aq) E = -1.812 V C. F2(g) + Br2(g) ® 2F-(aq) + 2Br-(aq) E = +1.812 V D. 2F (aq) + 2Br (aq) ® F2(g) + Br2(g) E = +1.812 V E. Br2(g) + 2F-(aq) ® 2Br-(aq) + F2(g) E = -1.812 V
31. The overall chemical reaction for the electrolytic decomposition of water is 2H2O ® O2(g) + 2H2(g) What is the reduction half-reaction? A. 2H+(aq) ® H2(g) + 2eB. H2(g) + 2OH-(aq) + 2e- ® 2H2O(l) C. 2H2O(l) + 2e- ® H2(g) + 2OH-(aq) D. O2(g) + 4H+(aq) + 4e- ® 2H2O(l) E. H2O(l) ® H2(g) + O2-(aq) + 2e-
32. For a spontaneous reaction, at constant pressure A. DG > 0 and E > 0. B. DG > 0 and E < 0. C. DG < 0 and E > 0. D. DG < 0 and E < 0. E. DG = 0 and E = 0.
33. What is the relationship between DG and E? A. DG = FE B. DG = -nFE C. D. E.
34. What is the relationship between E and the equilibrium constant, K? A. B. C. E = RT ln K D. E = -RT ln K E.
35. Calculate the equilibrium constant for the following reaction at 25C, 2IO3–(aq) + 5Hg(l) + 12H+(aq) ® I2(s) + 5Hg2+(aq) + 6H2O(l) given the following thermodynamic information. (R = 8.31 J/Kmol, F = 96480 C/mol e-) IO3–(aq) + 6H+(aq) + 5e– ® I2(s) + 3H2O(l) Hg2+(aq) + 2e– ® Hg(l)
E = +1.20 V E = +0.86 V
A. 3 ´ 10–58 B. 6 ´ 105 C. 3 ´ 1011 D. 6 ´ 1028 E. 3 ´ 1057 36. A voltaic cell is based upon the following overall reaction: 2Ag+(aq) + Sn(s) ® 2Ag(s) + Sn2+(aq) where E = +0.940 V. Calculate the equilibrium constant at 25C? (R = 8.31 J/Kmol, F = 96480 C/mol e-) A. 6.1 ´ 1012 B. 8.0 ´ 1015 C. 1.6 ´ 1023 D. 6.5 ´ 1031 E. 9.1 ´ 1043
37. A voltaic cell is based upon the half-reactions below. Pb2+(aq) + 2e- ® Pb(s) Cd2+(aq) + 2e- ® Cd(s)
E = -0.127 V E = -0.402 V
Calculate the equilibrium constant for the overall chemical reaction at 25C. (R = 8.31 J/Kmol, F = 96480 C/mol e-)
A. 1.3 ´ 102 B. 4.4 ´ 104 C. 2.0 ´ 109 D. 3.9 ´ 1018 E. 1.3 ´ 1021 38. Given the following standard reduction potentials, Pb2+(aq) + 2e– ® Pb(s) PbSO4(s) + 2e– ® Pb(s) + SO42–(aq)
E = –0.126 V E = –0.355 V
determine Ksp for PbSO4(s) at 25C. (R = 8.31 J/Kmol, F = 96480 C/mol e-)
A. 3.4 ´ 10–28 B. 1.8 ´ 10–8 C. 5.6 ´ 10–5 D. 5.6 ´ 107 E. 2.9 ´ 1037 39. Calculate DG for a voltaic cell with E = +0.24 V if the overall reaction involves a 3 electron reduction. (R = 8.31 J/Kmol, F = 96480 C/mol e-) A. -69 kJ B. -23 kJ C. -0.83 kJ D. + 220 kJ E. +580 kJ
40. Calculate DG for the disproportionation reaction of Cu+ at 25C, 2Cu+(aq) ® Cu2+(aq) + Cu(s) given the following thermodynamic information. (R = 8.31 J/Kmol, F = 96480 C/mol e-) Cu+(aq) + e– ® Cu(s) Cu2+(aq) + 2e– ® Cu(s)
A. –165 kJ B. –135 kJ C. –34.9 kJ D. +17.5 kJ E. +135 kJ
E = +0.518 V E = +0.337 V
41. What is the correct form of the Nernst equation? A. B. C. D. E.
42. In an electrochemical cell, if E > E then A. Q > 1. B. Q < 1. C. n = 1. D. F > 9.648 ´ 104. E. the system has reached equilibrium.
43. Calculate E for the following electrochemical cell at 25C, Cu | Cu2+ (0.100 M) || Zn2+ (0.0750 M) | Zn if E (Cu2+) = +0.339 V and E (Zn2+) = -0.762 V. (R = 8.31 J/Kmol, F = 96480 C/mol e-) A. -1.105 V B. -0.919 V C. -0.486 V D. -0.259 V E. +0.486 V
44. Calculate E for the following electrochemical cell at 25C, Pt | H2(g) (1.0 atm) | H+ (0.010 M) || Ag+ (0.020 M) | Ag if E (H+) = +0.000 V and E (Ag+) = 0.799 V. A. +0.275 V B. +0.799 V C. +0.817 V D. +0.911 V E. +1.01 V
45. Calculate E for the following electrochemical cell at 25C Pt | Fe3+(0.100 M), Fe2+(0.040 M) || Cl–(0.50 M) | AgCl | Ag given the following standard reduction potentials. AgCl(s) + e– ® Ag(s) + Cl–(aq) Fe3+(aq) + e– ® Fe2+(aq)
E = +0.222 V E = +0.771 V
A. –1.034 V B. –0.590 V C. –0.508 V D. –0.555 V E. +1.034 V 46. What concentration of Ag+ is present in the following electrochemical cell if E = -0.683 V and E (Ag+) = +0.799 V at 25C? Ag | Ag+ (? M) || H+ (1.0 M) | H2 (1.0 atm) | Pt A. 9.3 ´ 10-26 M B. 8.1 ´ 10-4 M C. 1.1 ´ 10-2 M D. 9.1 ´ 10-1 M E. 9.1 ´ 101 M
47. What is the pH of the solution at the cathode if E = –0.362 V for the following electrochemical cell at 25C? Pt | H2(1.0 atm) | H+(1.00 M) || H+(aq) | H2(1.0 atm) | Pt A. 1.77 B. 3.06 C. 6.11 D. 7.89 E. 12.23
48. Al3+ is reduced to Al(s) at an electrode surface. If a current of 25 amperes is maintained for 8.0 hours, what mass of aluminum is deposited on the electrode? Assume 100% current efficiency. A. 0.019 g B. 2.5 g C. 44 g D. 67 g E. 2.0 ´ 102 g
49. What charge, in coulombs, is required to deposit 1.5 g Mg(s) from a solution of Mg2+(aq)? (F = 96480 C/mol e-) A. 4.1 ´ 102 C B. 6.0 ´ 103 C C. 1.2 ´ 104 C D. 2.9 ´ 105 C E. 3.1 ´ 106 C
50. What charge (in Coulombs) is required to oxidize 1.0 g of silver to silver(I) ions? (F = 96480 C/mol e-) A. 1.0 ´ 10-5 C B. 1.1 ´ 10-3 C C. 3.1 ´ 10-1 C D. 6.1 ´ 101 C E. 8.9 ´ 102 C
51. An electrolytic cell is used to plate silver from a silver nitrate solution onto an electrode. How much time (in seconds) is required to deposit 8.0 g of silver if the current is kept constant at 5.0 amperes? Assume 100% current efficiency. (F = 96480 C/mol e-) A. 1.7 ´ 102 s B. 1.4 ´ 103 s C. 2.5 ´ 104 s D. 5.9 ´ 104 s E. 2.6 ´ 105 s
52. If current is passed in an electrolytic cell containing sodium chloride dissolved in water, what is the product at the cathode? A. H+ B. NaCl(s) C. H2(g) D. Cl2(g) E. O2(g)
53. One kind of battery used in watches contains mercury(II) oxide. As current flows, the mercury(II) oxide is reduced to mercury. HgO(s) + H2O(l) + 2e– ® Hg(l) + 2OH–(aq) If 2.3 ´ 10–5 amperes flows continuously for 1200 days, what mass of Hg(l) is produced? (F = 96480 C/mol e-) A. 2.5 g B. 5.0 g C. 9.9 g D. 13 g E. 15 g
54. Michael Faraday was A. a researcher, experimenting with electrolysis. B. a theoretician, who studied electrical charge. C. never formally educated past elementary school. D. both a and c. E. both b and c.
55. Fuel cells are being researched as a possible replacement for the internal combustion engine (ICE) in automobiles. What is/are the advantages of using a fuel cell/electric motor in place of the (ICE) in an automobile? A. H2 fuel cells only produce H2O vs. (ICE) produces NO(g) and CO2(g). B. Fuel cells are more efficient. C. Fuel cells are less expensive. D. both a and b. E. all of the above.
56. A baby's spoon with an area of 6.25cm2 is plated with silver from a silver nitrate solution using a current of 2.00 A for two hours and 25 minutes. What is the thickness of the silver plate formed if the current efficiency is 82%? (F = 96480 C/mol e-) (d Ag= 10.5 g/cm3) A. 0.300 cm B. 0.243 cm C. 0.366 cm D. 0.166 cm E. 5.34 ´ 10-2 cm
57. For the lead storage battery, the products of the discharge reaction are A. Pb at the anode, H2O at the cathode. B. H2SO4 at the anode, H2O at the cathode. C. PbSO4 at both the anode and cathode. D. Pb at the cathode and PbO2 at the anode. E. Pb at both the anode and cathode.
Chapter 17--Electrochemistry Key
1. If a reactant loses electrons it is A. reduced. B. oxidized. C. disproportionated. D. the cathode. E. a catalyst.
2. The anode is where A. electrolyte concentrations are lowest. B. electrolyte concentrations are highest. C. the Nernst potential is measured. D. anions are produced. E. oxidation occurs.
3. All of the following statements concerning voltaic cells are true EXCEPT A. a salt bridge allows cations and anions to move between the half-cells. B. electrons flow from the cathode to the anode in the external circuit. C. oxidation occurs at the anode. D. a voltaic cell can be used as a source of energy. E. a voltaic cell consists of two-half cells.
4. The following reaction occurs spontaneously. 2 Fe(s) + 3 Cl2(aq) ® 2 Fe3+(aq) + 6 Cl–(aq) Write the balanced reduction half-reaction. A. Fe(s) ® Fe3+(aq) + e– B. Fe(s) + 3 e– ® Fe3+(aq) C. Fe(s) + 3/2 Cl2(aq) ® FeCl3(aq) D. Cl2(aq) ® 2 Cl–(aq) + 2 e– E. Cl2(aq) + 2 e– ® 2 Cl–(aq)
5. The following reaction occurs spontaneously. 2Al(s) + 6H+(aq) ® 2Al3+(aq) + 3H2(g) Write the balanced oxidation half-reaction. A. Al(s) ® Al3+(aq) + 3eB. Al3+(aq) + 3e- ® Al(s) C. H2(g) ® 2H+(aq) + 2eD. 2H+(aq) + 2e- ® H2(g) E. 2H+(aq) ® H2(g) + 2e-
6. Given the following half-reactions, Pb(s) + 2I-(aq) ® PbI2(s) + 2eI2(aq) + 2e- ® 2I-(aq) write the balanced overall reaction. A. Pb(s) + I2(aq) + 2e- ® PbI (s) + I-(aq) B. Pb(s) + I2(aq) ® PbI2(s) C. Pb(s) + Cl2(aq) ® Pb2+(aq) + 2I-(aq) D. 2Pb(s) + 4e- ® 2PbI(s) + I2(aq) E. Pb(s) + e- + 2I-(aq) ® PbI2(aq)
7. Given the following half-reactions, 2IO3-(aq) + 12H+(aq) + 10e- ® I2(s) + 6H2O(l) Fe2+(aq) ® Fe3+(aq) + ewrite the balanced overall reaction. A. 2IO3-(aq) + Fe2+(aq) ® I2(s) + Fe3+(aq) B. 2IO3-(aq) + 10Fe2+(aq) ® I2(s) + 10Fe3+(aq) C. 2IO3-(aq) + 12H+(aq) + Fe2+(aq) ® FeI2(s) + 6H2O(l) D. 2IO3-(aq) + 12H+(aq) + Fe2+(aq) ® I2(s) + Fe3+(aq) + 6H2O(l) E. 2IO3-(aq) + 12H+(aq) + 10Fe2+(aq) ® I2(s) + 10Fe3+(aq) + 6H2O(l)
8. Write a balanced net ionic equation for the overall reaction represented by the cell notation below. Ag | AgI | I– || Fe2+, Fe3+ | Pt A. Fe3+(aq) + Ag(s) + I–(aq) ® Fe2+(aq) + AgI(s) B. Fe2+(aq) + Ag(s) + I–(aq) ® Fe3+(aq) + AgI(s) C. Fe3+(aq) + AgI(s) ® Fe2+(aq) + Ag(s) + I–(aq) D. Fe2+(aq) + AgI(s) ® Fe3+(aq) + Ag(s) + I–(aq) E. Fe3+(aq) + 2I–(aq) ® Fe2+(aq) + I2(s)
9. Write a balanced chemical equation for the overall reaction represented below. Pt | Fe2+, Fe3+ || Cl- | AgCl | Ag A. AgCl(s) + Fe3+(aq) + 2e- ® Ag(s) + Fe2+(aq) + Cl-(aq) B. AgCl(s) + Fe3+(aq) ® Ag(s) + Fe2+(aq) + Cl-(aq) C. AgCl(s) + Fe2+(aq) ® Ag(s) + Fe3+(aq) + Cl-(aq) D. Ag(s) + Fe3+(aq) + Cl-(aq) ® AgCl(s) + Fe2+(aq) E. Ag(s) + Fe2+(aq) + Cl-(aq) ® AgCl(s) + Fe3+(aq)
10. Write a balanced chemical equation for the overall reaction represented below. Cu | Cu2+ || Cl- | Hg2Cl2 | Hg A. 2Hg(l) + Cu2+(aq) + 2Cl-(aq) ® Hg2Cl2(s) + Cu(s) B. Hg2Cl2(s) + Cu(s) ® 2Hg(l) + Cu2+(aq) + 2Cl-(aq) C. Hg2Cl2(s) ® 2Hg(l) + Cl2(aq) D. Hg2Cl2(s) + Cu(s) ® 2Hg(l) + Cu2+(aq) + Cl2(aq) E. 2Hg(l) + Cl2(aq) ® Hg2Cl2(s)
11. What is the correct cell notation for the reaction below? Cd2+(aq) + Ni(s) ® Cd(s) + Ni2+(aq) A. Cd2+ | Cd || Ni | Ni2+ B. Cd | Cd2+ || Ni2+ | Ni C. Ni | Ni2+ || Cd2+ | Cd D. Ni | Cd2+ || Ni2+ | Cd E. Cd2+ | Cd || Ni | Ni2+
12. What is a correct cell notation for a voltaic cell based on the reaction below? Cu2+(aq) + Pb(s) + SO42–(aq) ® Cu(s) + PbSO4(s) A. Pb | PbSO4 || Cu2+ || Cu B. Cu | Cu2+ || SO42– | PbSO4 | Pb C. Cu | Cu2+, SO42– | PbSO4 | Pb D. Cu | Cu2+, SO42– || PbSO4 || Pb E. Pb | PbSO4 | SO42– || Cu2+ || Cu
13. Which species is the best oxidizing agent? A. bromine B. sodium ion C. potassium metal D. fluoride ion E. carbon dioxide
14. Which species is the best reducing agent? A. sodium ion B. sulfide ion C. chlorine D. sodium metal E. silver metal
15. Use the standard reduction potentials below to determine which element or ion is the best reducing agent. NO3–(aq) + 4H+(aq) + 3e– ® NO(g) + 2H2O(l) Pd2+(aq) + 2e– ® Pd(s) 2H+(aq) + 2e– ® H2(g)
E = +0.955 V E = +0.90 V E = 0.00 V
A. Pd2+(aq) B. Pd(s) C. H+(aq) D. H2(g) E. NO3–(aq)
16. The standard reduction potential, oxidation potential. What is the value of
, of Zn2+(aq) is -0.762 V. Write the equation for the standard ?
A. Zn2+(aq) + 2e- ® Zn(s)
= -0.762 V
B. Zn2+(aq) ® Zn(s) + 2e-
= +0.762 V
C. Zn2+(aq) ® Zn(s) + 2e-
= -0.762 V
D. Zn(s) + 2e- ® Zn2+(aq)
= -0.762 V
E. Zn(s) ® Zn2+(aq) + 2e-
= +0.762 V
17. The standard oxidation potential, , of Cu(s) is -0.339 V. The oxidation product is the copper(II) ion. Write the equation for the standard reduction potential. What is the value of the standard reduction potential, ? A. Cu2+(aq) ® Cu(s) + 2e-
= -0.339 V
B. Cu2+(aq) + 2e- ® Cu(s)
= +0.339 V
C. Cu(s) + 2e- ® Cu2+(aq)
= -0.339 V
D. Cu(s) ® Cu2+(aq) + 2e-
= +0.339 V
E. Cu(s) + 2e- ® Cu2+(aq)
= +0.339 V
18. Which of the following species are likely to behave as oxidizing agents: Li(s), H2(g), MnO4–(aq), and Cl–(aq)? A. Li(s) only B. MnO4–(aq) only C. H2(g) and Cl–(aq) D. Li(s) and MnO4–(aq) E. Cl–(aq) only
19. Which of the following species are reducing agents: K, H+, MnO4-, Cl2, Sn4+? A. K only B. K and Cl2 C. H+ and Sn4+ D. MnO4- and Cl2 E. Cl2 and Sn4+
20. Write the overall chemical reaction and calculate E for the following reduction and oxidation half-reactions. 2H+(aq) + 2e- ® H2(g) = 0.000 V Zn(s) ® Zn2+(aq) + 2e= +0.762 V
A. Zn2+(aq) + H2(g) ® Zn(s) + 2H+(aq) B. Zn2+(aq) + H2(g) ® Zn(s) + 2H+(aq) C. Zn(s) + 2H+(aq) ® Zn2+(aq) + H2(g) D. Zn(s) + 2H+(aq) ® Zn2+(aq) + H2(g) E. Zn(s) + 2H+(aq) ® Zn2+(aq) + H2(g)
E = +0.381 V E = -0.762 V E = +0.762 V E = -0.762 V E = -0.381 V
21. Given the following two half-reactions, write the overall balanced reaction in the direction in which it is spontaneous and calculate the standard cell potential. Ga3+(aq) + 3e– ® Ga(s) = –0.53 V Sn4+(aq) + 2e– ® Sn2+(aq) = +0.15 V
A. 2Ga3+(aq) + 3Sn2+(aq) ® 2Ga(s) + 3Sn4+(aq) B. 3Ga3+(aq) + 2Sn2+(aq) ® 3Ga(s) + 2Sn4+(aq) C. 2Ga(s) + 3Sn4+(aq) ® 2Ga3+(aq) + 3Sn2+(aq) D. 3Ga(s) + 2Sn4+(aq) ® 3Ga3+(aq) + 2Sn2+(aq) E. 2Ga(s) + 3Sn4+(aq) ® 2Ga3+(aq) + 3Sn2+(aq)
E = +0.68 V E = –1.89 V E = +0.68 V E = +0.38 V E = +1.89 V
22. Using the reduction half-reactions, write an overall reaction in the direction that is spontaneous. Calculate E. AgI(s) + e- ® Ag(s) + I-(aq) = -0.152 V Co3+(aq) + e- ® Co2+(aq) = +1.953 V
A. Co2+(aq) + AgI(s) ® Co3+(aq) + Ag(s) + I-(aq) B. Co2+(aq) + Ag(s) + I-(aq) ® Co3+(aq) + AgI(s) C. Co3+(aq) + Ag(s) + I-(aq) ® Co2+(aq) + AgI(s) D. Co3+(aq) + Ag(s) + I-(aq) ® Co2+(aq) + AgI(s) E. Co3+(aq) + AgI(s) ® Co2+(aq) + Ag(s) + I-(aq)
E = -2.105 V E = -1.801 V E = +1.801 V E = +2.105 V E = -2.105 V
23. Using the reduction half-reactions, write an overall reaction in the direction that it is spontaneous. Calculate E. MnO4-(aq) + 8H+(aq) + 5e- ® 2Mn2+(aq) + 4H2O(l) = +1.512 V Fe3+(aq) + e- ® Fe2+(aq) = +0.769 V
A. 5Fe3+(aq) + Mn2+(aq) + 4H2O(l) ® 5Fe2+(aq) + MnO4-(aq) + 8H+(aq) B. 5Fe2+(aq) + Mn2+(aq) + 4H2O(l) ® 5Fe3+(aq) + MnO4-(aq) + 8H+(aq) C. MnO4-(aq) + 5Fe2+(aq) ® Mn2+(aq) + 5Fe3+(aq) E = -2.281 V D. MnO4-(aq) + 8H+(aq) + 5Fe2+(aq) ® Mn2+(aq) + 5Fe3+(aq) + 4H2O(l) E. MnO4-(aq) + 8H+(aq) + 5Fe3+(aq) ® Mn2+(aq) + 5Fe2+(aq) + 4H2O(l)
E = +2.333 V E = -0.743 V E = +0.743 V E = +2.281 V
24. The cell voltage, E, is ____ for a reaction taking place in a voltaic cell. A. reductive B. oxidative C. negative D. zero E. positive
25. Which of the following half-reactions has been assigned a standard reduction potential of 0.00 V? A. 2H+(aq) + 2e- ® H2(g) B. Ag+(aq) + e- ® Ag(s) C. Hg2Cl2(s) + 2e- ® 2Hg(l) + 2Cl-(aq) D. Au3+(aq) + 3e- ® Au(s) E. Cu2+(aq) + 2e- ® Cu(s)
26. Calculate E for an electrochemical cell based on the following overall reaction: 2H+(aq) + Cd(s) ® H2(g) + Cd2+(aq) if (H+) = 0.000 V and A. -0.402 V B. -0.201 V C. 0.000 V D. +0.201 V E. +0.402 V
(Cd2+) = -0.402 V.
27. Calculate E for the electrochemical cell below, Ag | AgCl| Cl– || Cu2+ | Cu given the following standard reduction potentials. Cu2+(aq) + 2e– ® Cu(s) AgCl(s) + e– ® Ag(s) + Cl–(aq)
E = +0.337 V E = +0.222 V
A. –0.115 V B. –0.107 V C. +0.115 V D. +0.452 V E. +0.559 V 28. Use the standard reduction potentials below to determine which of the following metals can be oxidized by 1.0 M HCl: Al, Zn, Sn, Cu. Al3+(aq) + 3e- ® Al(s) = -1.68 V Zn2+(aq) + 2e- ® Zn(s) = -0.762 V Sn2+(aq) + 2e- ® Sn(s) = -0.141 V Cu2+(aq) + 2e- ® Cu(s) = +0.339 V
A. Al only B. Cu only C. Al and Cu D. Zn and Sn E. Al, Zn, and Sn 29. Calculate E for the electrochemical cell below, Pb |PbCl2 | Cl– || Fe3+, Fe2+ | Pt given the following reduction half-reactions. Pb2+(aq) + 2e– ® Pb(s) PbCl2(s) + 2e– ® Pb(s) + 2Cl–(aq) Fe3+(aq) + e– ® Fe2+(aq) Fe2+(aq) + e– ® Fe(s)
A. –0.504 V B. –0.062 V C. +0.504 V D. +1.038 V E. +1.604 V
E = –0.126 V E = –0.267 V E = +0.771 V E = –0.44 V
30. The standard reduction potentials for F2(g) and Br2(l) are +2.889 V and 1.077 V, respectively. Write the chemical equation and calculate E for a voltaic cell based on fluorine and bromine. A. F2(g) + 2Br-(aq) ® 2F-(aq) + Br2(g) E = +1.812 V B. F2(g) + Br2(g) ® 2F-(aq) + 2Br-(aq) E = -1.812 V C. F2(g) + Br2(g) ® 2F-(aq) + 2Br-(aq) E = +1.812 V D. 2F (aq) + 2Br (aq) ® F2(g) + Br2(g) E = +1.812 V E. Br2(g) + 2F-(aq) ® 2Br-(aq) + F2(g) E = -1.812 V
31. The overall chemical reaction for the electrolytic decomposition of water is 2H2O ® O2(g) + 2H2(g) What is the reduction half-reaction? A. 2H+(aq) ® H2(g) + 2eB. H2(g) + 2OH-(aq) + 2e- ® 2H2O(l) C. 2H2O(l) + 2e- ® H2(g) + 2OH-(aq) D. O2(g) + 4H+(aq) + 4e- ® 2H2O(l) E. H2O(l) ® H2(g) + O2-(aq) + 2e-
32. For a spontaneous reaction, at constant pressure A. DG > 0 and E > 0. B. DG > 0 and E < 0. C. DG < 0 and E > 0. D. DG < 0 and E < 0. E. DG = 0 and E = 0.
33. What is the relationship between DG and E? A. DG = FE B. DG = -nFE C. D. E.
34. What is the relationship between E and the equilibrium constant, K? A. B. C. E = RT ln K D. E = -RT ln K E.
35. Calculate the equilibrium constant for the following reaction at 25C, 2IO3–(aq) + 5Hg(l) + 12H+(aq) ® I2(s) + 5Hg2+(aq) + 6H2O(l) given the following thermodynamic information. (R = 8.31 J/Kmol, F = 96480 C/mol e-) IO3–(aq) + 6H+(aq) + 5e– ® I2(s) + 3H2O(l) Hg2+(aq) + 2e– ® Hg(l)
E = +1.20 V E = +0.86 V
A. 3 ´ 10–58 B. 6 ´ 105 C. 3 ´ 1011 D. 6 ´ 1028 E. 3 ´ 1057 36. A voltaic cell is based upon the following overall reaction: 2Ag+(aq) + Sn(s) ® 2Ag(s) + Sn2+(aq) where E = +0.940 V. Calculate the equilibrium constant at 25C? (R = 8.31 J/Kmol, F = 96480 C/mol e-) A. 6.1 ´ 1012 B. 8.0 ´ 1015 C. 1.6 ´ 1023 D. 6.5 ´ 1031 E. 9.1 ´ 1043
37. A voltaic cell is based upon the half-reactions below. Pb2+(aq) + 2e- ® Pb(s) Cd2+(aq) + 2e- ® Cd(s)
E = -0.127 V E = -0.402 V
Calculate the equilibrium constant for the overall chemical reaction at 25C. (R = 8.31 J/Kmol, F = 96480 C/mol e-)
A. 1.3 ´ 102 B. 4.4 ´ 104 C. 2.0 ´ 109 D. 3.9 ´ 1018 E. 1.3 ´ 1021 38. Given the following standard reduction potentials, Pb2+(aq) + 2e– ® Pb(s) PbSO4(s) + 2e– ® Pb(s) + SO42–(aq)
E = –0.126 V E = –0.355 V
determine Ksp for PbSO4(s) at 25C. (R = 8.31 J/Kmol, F = 96480 C/mol e-)
A. 3.4 ´ 10–28 B. 1.8 ´ 10–8 C. 5.6 ´ 10–5 D. 5.6 ´ 107 E. 2.9 ´ 1037 39. Calculate DG for a voltaic cell with E = +0.24 V if the overall reaction involves a 3 electron reduction. (R = 8.31 J/Kmol, F = 96480 C/mol e-) A. -69 kJ B. -23 kJ C. -0.83 kJ D. + 220 kJ E. +580 kJ
40. Calculate DG for the disproportionation reaction of Cu+ at 25C, 2Cu+(aq) ® Cu2+(aq) + Cu(s) given the following thermodynamic information. (R = 8.31 J/Kmol, F = 96480 C/mol e-) Cu+(aq) + e– ® Cu(s) Cu2+(aq) + 2e– ® Cu(s)
A. –165 kJ B. –135 kJ C. –34.9 kJ D. +17.5 kJ E. +135 kJ
E = +0.518 V E = +0.337 V
41. What is the correct form of the Nernst equation? A. B. C. D. E.
42. In an electrochemical cell, if E > E then A. Q > 1. B. Q < 1. C. n = 1. D. F > 9.648 ´ 104. E. the system has reached equilibrium.
43. Calculate E for the following electrochemical cell at 25C, Cu | Cu2+ (0.100 M) || Zn2+ (0.0750 M) | Zn if E (Cu2+) = +0.339 V and E (Zn2+) = -0.762 V. (R = 8.31 J/Kmol, F = 96480 C/mol e-) A. -1.105 V B. -0.919 V C. -0.486 V D. -0.259 V E. +0.486 V
44. Calculate E for the following electrochemical cell at 25C, Pt | H2(g) (1.0 atm) | H+ (0.010 M) || Ag+ (0.020 M) | Ag if E (H+) = +0.000 V and E (Ag+) = 0.799 V. A. +0.275 V B. +0.799 V C. +0.817 V D. +0.911 V E. +1.01 V
45. Calculate E for the following electrochemical cell at 25C Pt | Fe3+(0.100 M), Fe2+(0.040 M) || Cl–(0.50 M) | AgCl | Ag given the following standard reduction potentials. AgCl(s) + e– ® Ag(s) + Cl–(aq) Fe3+(aq) + e– ® Fe2+(aq)
E = +0.222 V E = +0.771 V
A. –1.034 V B. –0.590 V C. –0.508 V D. –0.555 V E. +1.034 V 46. What concentration of Ag+ is present in the following electrochemical cell if E = -0.683 V and E (Ag+) = +0.799 V at 25C? Ag | Ag+ (? M) || H+ (1.0 M) | H2 (1.0 atm) | Pt A. 9.3 ´ 10-26 M B. 8.1 ´ 10-4 M C. 1.1 ´ 10-2 M D. 9.1 ´ 10-1 M E. 9.1 ´ 101 M
47. What is the pH of the solution at the cathode if E = –0.362 V for the following electrochemical cell at 25C? Pt | H2(1.0 atm) | H+(1.00 M) || H+(aq) | H2(1.0 atm) | Pt A. 1.77 B. 3.06 C. 6.11 D. 7.89 E. 12.23
48. Al3+ is reduced to Al(s) at an electrode surface. If a current of 25 amperes is maintained for 8.0 hours, what mass of aluminum is deposited on the electrode? Assume 100% current efficiency. A. 0.019 g B. 2.5 g C. 44 g D. 67 g E. 2.0 ´ 102 g
49. What charge, in coulombs, is required to deposit 1.5 g Mg(s) from a solution of Mg2+(aq)? (F = 96480 C/mol e-) A. 4.1 ´ 102 C B. 6.0 ´ 103 C C. 1.2 ´ 104 C D. 2.9 ´ 105 C E. 3.1 ´ 106 C
50. What charge (in Coulombs) is required to oxidize 1.0 g of silver to silver(I) ions? (F = 96480 C/mol e-) A. 1.0 ´ 10-5 C B. 1.1 ´ 10-3 C C. 3.1 ´ 10-1 C D. 6.1 ´ 101 C E. 8.9 ´ 102 C
51. An electrolytic cell is used to plate silver from a silver nitrate solution onto an electrode. How much time (in seconds) is required to deposit 8.0 g of silver if the current is kept constant at 5.0 amperes? Assume 100% current efficiency. (F = 96480 C/mol e-) A. 1.7 ´ 102 s B. 1.4 ´ 103 s C. 2.5 ´ 104 s D. 5.9 ´ 104 s E. 2.6 ´ 105 s
52. If current is passed in an electrolytic cell containing sodium chloride dissolved in water, what is the product at the cathode? A. H+ B. NaCl(s) C. H2(g) D. Cl2(g) E. O2(g)
53. One kind of battery used in watches contains mercury(II) oxide. As current flows, the mercury(II) oxide is reduced to mercury. HgO(s) + H2O(l) + 2e– ® Hg(l) + 2OH–(aq) If 2.3 ´ 10–5 amperes flows continuously for 1200 days, what mass of Hg(l) is produced? (F = 96480 C/mol e-) A. 2.5 g B. 5.0 g C. 9.9 g D. 13 g E. 15 g
54. Michael Faraday was A. a researcher, experimenting with electrolysis. B. a theoretician, who studied electrical charge. C. never formally educated past elementary school. D. both a and c. E. both b and c.
55. Fuel cells are being researched as a possible replacement for the internal combustion engine (ICE) in automobiles. What is/are the advantages of using a fuel cell/electric motor in place of the (ICE) in an automobile? A. H2 fuel cells only produce H2O vs. (ICE) produces NO(g) and CO2(g). B. Fuel cells are more efficient. C. Fuel cells are less expensive. D. both a and b. E. all of the above.
56. A baby's spoon with an area of 6.25cm2 is plated with silver from a silver nitrate solution using a current of 2.00 A for two hours and 25 minutes. What is the thickness of the silver plate formed if the current efficiency is 82%? (F = 96480 C/mol e-) (d Ag= 10.5 g/cm3) A. 0.300 cm B. 0.243 cm C. 0.366 cm D. 0.166 cm E. 5.34 ´ 10-2 cm
57. For the lead storage battery, the products of the discharge reaction are A. Pb at the anode, H2O at the cathode. B. H2SO4 at the anode, H2O at the cathode. C. PbSO4 at both the anode and cathode. D. Pb at the cathode and PbO2 at the anode. E. Pb at both the anode and cathode.
Chapter 18--Nuclear Reactions 1. All of the following statements concerning nuclei are true EXCEPT A. only hydrogen-1 and helium-3 have more protons than neutrons. B. from He to Ca, stable nuclei have roughly equal numbers of protons and neutrons. C. elements with odd atomic numbers have more stable isotopes than do those with even atomic numbers. D. the neutron to proton ratio in stable nuclei increases as mass increases. E. Elements with an atomic number greater than 83 are always radioactive.
2. A plot of the number of neutrons versus the number of protons in nuclei shows a narrow band of stable isotopes. By what method(s) can isotopes with a low proton-neutron ratio (i.e., ones that fall below the band of stability) decay to form elements that are more stable? A. positron emission or electron capture. B. positron emission or neutron capture. C. beta emission or electron capture. D. gamma ray emission or beta emission. E. neutron capture or alpha emission.
3. Which of the following reactions is an example of alpha particle emission? A. B. C. D. E.
4. Which of the following reactions is an example of beta particle emission? A. B. C. D. E.
5. Which of the following reactions is an example of K-electron capture? A. B. C. D. E.
6. Which of the following reactions is an example of positron emission? A. B. C. D. E.
7. Which of the following types of radiation will pass through a piece of paper, but will be stopped by 0.5 cm of lead? A. a B. b C. D. g E. All of the above will pass through 0.5 cm of lead.
8. If a nucleus emits an alpha particle, its mass number A. decreases by 4. B. decreases by 2. C. remains the same. D. increases by 1. E. decreases by 2.
9. If a nucleus captures an electron, its atomic number A. decreases by 2 B. decreases by 1. C. remains the same. D. increases by 1. E. increases by 2.
10. If a nucleus emits a beta particle, its atomic number A. decreases by 2. B. decreases by 1. C. remains the same. D. increases by 1. E. increases by 2.
11. If a nucleus decays by successive a, b, b particle emissions, its atomic number will A. decrease by four. B. decrease by two. C. increase by four. D. increase by two. E. be unchanged.
12. All of the following statements are true EXCEPT A. gamma emission can accompany nuclear decay, such as beta emission. B. gamma rays can be used to extend the shelf life of food products. C. gamma rays have the same properties as electrons. D. gamma emission consists of high energy photons. E. gamma rays move at the speed of light.
13. What element is produced by the alpha decay of ? A. B. C. D. E.
14. What nucleus decays by beta emission to produce antimony-121? A. B. C. D. E.
15. What nucleus decays by successive b, b, a emissions to produce uranium-236? A. B. C. D. E.
16. Complete the following fusion reaction. A. B. C. D. E.
17. Complete the following nuclear decay reaction. A. B. C. D. E.
18. Complete the following nuclear reaction. A. B. C. D. E.
19. Write the balanced nuclear equation for the formation of from beta emission. A. B. C. D. E.
20. Write a balanced nuclear equation for the positron emission reaction undergone by . A. B. C. D. E.
21. undergoes electron capture. Write a balanced nuclear equation for this process. A. B. C. D. E.
22. Complete the following nuclear fission reaction. A. B. C. D. E.
23. Complete the following nuclear fission reaction. A. B. C. D. E.
24. The following nuclear reaction is an example of ____. A. fusion B. fission C. gamma radiation emission D. beta emission E. hydrogen combustion
25. The following reaction is an example of ____. A. fusion B. fission C. gamma radiation emission D. beta emission E. neutron bombardment
26. What is the relationship between the rate constant and the half-life for nuclear decay. A. rate = kt1/2 B. C. k = ln (0.693t1/2) D. E. k = 0.693t1/2
27. Uranium-235 has a half-life of 7.04 ´ 108 years. How many years will it take for 99.9% of a U-235 sample to decay? A. 7.0 ´ 105 yr B. 1.0 ´ 106 yr C. 4.7 ´ 109 yr D. 4.9 ´ 109 yr E. 7.0 ´ 109 yr
28. How long does it take for 105 mg of tritium, , to decay to 35.0 mg? The half-life of tritium is 12.3 yr. A. 12.3 yr B. 14.2 yr C. 19.5 yr D. 24.6 yr E. 36.9 yr
29. The half-life of mercury-203 is 46.6 days. What percentage of mercury-203 remains in a sample after 365 days? A. 3.72 ´ 10-4 % B. 0.439% C. 1.49% D. 5.43% E. 12.8%
30. How many atoms decay in 12 hours in a 3.5 ´ 10-8 Ci source? (1 Ci = 3.700 ´ 1010 atom/s) A. 4.1 ´ 10-14 atoms B. 1.6 ´ 104 atoms C. 5.0 ´ 104 atoms D. 5.6 ´ 107 atoms E. 4.6 ´ 1022 atoms
31. The half-life of americium-241 is 457 years. What is the activity in curies of a 0.010 g sample of Am-241? (1 Ci = 3.700 ´ 1010 atom/s) A. 0.032 Ci B. 1.0 Ci C. 1.0 ´ 106 Ci D. 1.0 ´ 109 Ci E. 3.79 ´ 1016 Ci
32. A 1.00 g sample of decays to 0.87 g in 1.9 hr. What is the half-life of ? A. 0.073 hr B. 0.11 hr C. 4.5 hr D. 9.5 hr E. 14 hr
33. decays by beta emission. It has a decay constant of 7.64 ´ 10-10 s-1. How many beta particles are emitted per second from 1.00 ´ 10-2 moles of ? A. 7.64 ´ 10-12 s-1 B. 7.64 ´ 10-10 s-1 C. 3.12 ´ 105 s-1 D. 1.45 ´ 1010 s-1 E. 4.60 ´ 1012 s-1
34. decays by alpha emission. The isotope's half-life is 2.4 ´ 104 yr. How many alpha particles are emitted per second by 1.0 ´ 10-4 mol . A. 9.2 ´ 10-13 s-1 B. 3.0 ´ 10-9 s-1 C. 5.5 ´ 107 s-1 D. 7.6 ´ 1011 s-1 E. 4.8 ´ 1016 s-1
35. The rate constant for the decay of copper-62 is 9.8 minutes-1. What is the half-life of this isotope? A. 0.071 m. B. 1.6 m. C. 14 m. D. 0.10 m. E. 9.8 m.
36. The half-life of is 9.5 ´ 1012 s. How many atoms of are present in a sample with an activity of 3.9 ´ 109 atoms/s? A. 1.9 ´ 10-23 atoms B. 4.1 ´ 10-4 atoms C. 2.4 ´ 103 atoms D. 9.1 ´ 1016 atoms E. 5.3 ´ 1022 atoms
37. The decay constant, k, for is 1.37 ´ 10-11 s-1. What mass of has an activity of 1.06 ´ 109 atoms/s? A. 6.60 ´ 10-6 g B. 1.49 ´ 10-3 g C. 2.90 ´ 10-2 g D. 3.44 ´ 101 g E. 8.97 ´ 102 g
38. is a naturally occurring radioisotope with a half-life of 7.1 ´ 108 yr. What percentage of that was present during the formation of Earth still remains? Assume that the Earth is 4.5 ´ 109 yr old and that no is created by decomposition of other elements. A. 0.16% B. 0.81% C. 1.2% D. 4.4% E. 6.3%
39. is used in many home smoke alarms. It has a decay half-life of 457 years. How long does it take for the americium in a smoke alarm to decay to 6.25% of its original mass? A. 870 yr B. 1300 yr C. 1426 yr D. 1828 yr E. 2600 yr
40. All of the following statements concerning radiocarbon dating are true EXCEPT A. radiocarbon dating is used to determine the age of meteorites. B. the decay of is used for radiocarbon dating. C. radioactive carbon is produced by the reaction of cosmic radiation and nitrogen-14. D. radioactive carbon decays into a nitrogen-14 atom and a beta particle. E. radiocarbon dating is limited to objects less than 50,000 years old.
41. A 1.00 g sample of carbon taken from a fossilized plant has an activity of 0.19 atoms/s. What mass of is present? The decay constant of is 3.8 ´ 10-12 s-1. A. 4.6 ´ 10-34 g B. 5.0 ´ 10-14 g C. 1.2 ´ 10-12 g D. 6.1 ´ 10-12 g E. 3.2 ´ 10-11 g
42. A 1.00 g sample of carbon taken from a living plant has an activity of 13.6 atoms/min. If 1.00 g of an organic sample is found to have an activity of 6.27 atoms/min, what is the age of the sample? The decay constant of is 1.21 ´ 10-4 yr-1. A. 2.64 ´ 103 yr B. 5.28 ´ 103 yr C. 6.40 ´ 103 yr D. 8.18 ´ 103 yr E. 1.07 ´ 104 yr
43. Carbon dating can be applied to organic materials up to 5.0 ´ 104 years old. What percentage of the original is present in a sample after this much time? The half-life of is 5730 yr. A. 0.0011% B. 0.24% C. 0.89% D. 6.0% E. 8.7%
44. An extremely sensitive method has recently been developed for measuring concentrations in organic materials. What is this method? A. acid titration B. nuclear magnetic resonance C. gas chromatography D. atomic absorption spectrometry E. mass spectrometry
45. The energy change for a nuclear reaction can be calculated from which of the following equations? A. B. DE = hv C. DE = cDm2 D. DE = c2Dm E.
46. The point of maximum stability in the binding energy curve occurs in the vicinity of which one of the following isotopes? A. B. C. D. E.
47. All of the following statements are true EXCEPT A. a nucleus weighs less than the protons and neutrons from which it is composed. B. the difference in mass between a nucleus and its individual protons and neutrons is called the mass defect. C. the process in which a nucleus decomposes to form two or more lighter nuclei is called nuclear fission. D. nuclear fission evolves more energy than ordinary chemical reactions, such as the combustion of gasoline. E. nuclear fission evolves more energy than nuclear fusion.
48. All of the following statements are true EXCEPT A. for a nuclear fission chain reaction to occur, one neutron must be produced for each one consumed. B. in a light water nuclear reactor, steam drives the turbines used to generate electricity. C. in a fission reactor, reacts with a neutron to produce more neutrons and products of lower mass. D. many of the fission products of are radioactive. E. one function of the water in a light water reactor is to slow down neutrons given off by fission.
49. The molar nuclear mass of fluorine-19 is 18.99840 g/mol. The molar mass of a proton is 1.007825 g/mol. The molar mass of a neutron is 1.008665 g/mol. Calculate the binding energy (in kJ/mol) of F-19. (c = 2.998 ´ 108 m/s) A. 6.753 ´ 109 kJ/mol B. 7.131 ´ 109 kJ/mol C. 1.426 ´ 1010 kJ/mol D. 8.609 ´ 1011 kJ/mol E. 8.538 ´ 1011 kJ/mol
50. Calculate the energy change per mole for the following reaction? The mass of is 2.01355 amu and the mass of is 4.00150 amu. A. 2.30 ´ 109 kJ/mol B. 9.00 ´ 109 kJ/mol C. 1.79 ´ 1011 kJ/mol D. 3.52 ´ 1012 kJ/mol E. 2.84 ´ 1013 kJ/mol
51. Calculate the energy released (per mole of deuterium consumed) for the following fusion reaction, given the following molar masses of nucleons and nuclei. (c = 2.998 ´ 108 m/s)
particle proton neutron deuterium tritium helium-4
mass (g/mol) 1.007825 1.008665 2.0140 3.01605 4.00260
A. 5.63 ´ 103 kJ/mol B. 1.69 ´ 1012 kJ/mol C. 4.62 ´ 1010 kJ/mol D. 8.44 ´ 108 kJ/mol E. 1.69 ´ 109 kJ/mol 52. The following fusion reaction releases 2.16 ´ 109 kJ per mole of deuterium reacted. How much mass is lost in the reaction per mole of deuterium reacted? A. 1.2 ´ 10-5 g B. 2.4 ´ 10-5 g C. 1.2 ´ 10-2 g D. 2.4 ´ 10-2 g E. 1.2 ´ 10-1 g
53. decays by positron emission. Calculate the energy released or gained in this reaction using the information below.
mass (amu) 0.00055 1.00867 1.00728 11.00656 11.00814
A. 1.03 ´ 10-3 kJ/mol evolved in reaction B. 9.27 ´ 107 kJ/mol evolved in reaction C. 1.42 ´ 108 kJ/mol evolved in reaction D. 9.12 ´ 109 kJ/mol absorbed in reaction E. 8.74 ´ 1013 kJ/mol absorbed in reaction
54. All of the following statements are true EXCEPT A. the Sun's energy comes primarily from the fusion of hydrogen. B. fusion reactions have high activation barriers due to repulsion of nuclei. C. fission reactions have relatively low activation barriers because uncharged neutrons used to initiate the reactions are not repelled by nuclei. D. magnetic fields can be used to confine nuclei during a fusion reaction. E. many fission reactors have been replaced by fusion reactors, which produce more energy and fewer waste products.
55. Human exposure to radiation comes primarily from A. medical x-rays. B. sitting too close to the television. C. nuclear fallout from bomb tests. D. radon gas. E. using cellular telephones.
56. Marie Curie was the first person in history to A. discover radioactivity. B. die from leukemia. C. win two Nobel Prizes. D. name an element. E. discover the neutron.
57. Most of the sun's energy is produced primarily from A. nuclear fission of 235U. B. nuclear fusion of 235U. C. nuclear fission of 1H. D. nuclear fusion of 1H. E. nuclear fusion of 2H.
58. Most of the nuclear power on earth is produced primarily from A. nuclear fission of 235U. B. nuclear fusion of 235U. C. nuclear fission of 1H. D. nuclear fusion of 1H. E. nuclear fusion of 2H.
59. Radioactive elements have been used A. to determine the age of artifacts. B. to treat food to reduce spoilage. C. in smoke detectors. D. both a and b E. all of the above
Chapter 18--Nuclear Reactions Key
1. All of the following statements concerning nuclei are true EXCEPT A. only hydrogen-1 and helium-3 have more protons than neutrons. B. from He to Ca, stable nuclei have roughly equal numbers of protons and neutrons. C. elements with odd atomic numbers have more stable isotopes than do those with even atomic numbers. D. the neutron to proton ratio in stable nuclei increases as mass increases. E. Elements with an atomic number greater than 83 are always radioactive.
2. A plot of the number of neutrons versus the number of protons in nuclei shows a narrow band of stable isotopes. By what method(s) can isotopes with a low proton-neutron ratio (i.e., ones that fall below the band of stability) decay to form elements that are more stable? A. positron emission or electron capture. B. positron emission or neutron capture. C. beta emission or electron capture. D. gamma ray emission or beta emission. E. neutron capture or alpha emission.
3. Which of the following reactions is an example of alpha particle emission? A. B. C. D. E.
4. Which of the following reactions is an example of beta particle emission? A. B. C. D. E.
5. Which of the following reactions is an example of K-electron capture? A. B. C. D. E.
6. Which of the following reactions is an example of positron emission? A. B. C. D. E.
7. Which of the following types of radiation will pass through a piece of paper, but will be stopped by 0.5 cm of lead? A. a B. b C. D. g E. All of the above will pass through 0.5 cm of lead.
8. If a nucleus emits an alpha particle, its mass number A. decreases by 4. B. decreases by 2. C. remains the same. D. increases by 1. E. decreases by 2.
9. If a nucleus captures an electron, its atomic number A. decreases by 2 B. decreases by 1. C. remains the same. D. increases by 1. E. increases by 2.
10. If a nucleus emits a beta particle, its atomic number A. decreases by 2. B. decreases by 1. C. remains the same. D. increases by 1. E. increases by 2.
11. If a nucleus decays by successive a, b, b particle emissions, its atomic number will A. decrease by four. B. decrease by two. C. increase by four. D. increase by two. E. be unchanged.
12. All of the following statements are true EXCEPT A. gamma emission can accompany nuclear decay, such as beta emission. B. gamma rays can be used to extend the shelf life of food products. C. gamma rays have the same properties as electrons. D. gamma emission consists of high energy photons. E. gamma rays move at the speed of light.
13. What element is produced by the alpha decay of ? A. B. C. D. E.
14. What nucleus decays by beta emission to produce antimony-121? A. B. C. D. E.
15. What nucleus decays by successive b, b, a emissions to produce uranium-236? A. B. C. D. E.
16. Complete the following fusion reaction. A. B. C. D. E.
17. Complete the following nuclear decay reaction. A. B. C. D. E.
18. Complete the following nuclear reaction. A. B. C. D. E.
19. Write the balanced nuclear equation for the formation of from beta emission. A. B. C. D. E.
20. Write a balanced nuclear equation for the positron emission reaction undergone by . A. B. C. D. E.
21. undergoes electron capture. Write a balanced nuclear equation for this process. A. B. C. D. E.
22. Complete the following nuclear fission reaction. A. B. C. D. E.
23. Complete the following nuclear fission reaction. A. B. C. D. E.
24. The following nuclear reaction is an example of ____. A. fusion B. fission C. gamma radiation emission D. beta emission E. hydrogen combustion
25. The following reaction is an example of ____. A. fusion B. fission C. gamma radiation emission D. beta emission E. neutron bombardment
26. What is the relationship between the rate constant and the half-life for nuclear decay. A. rate = kt1/2 B. C. k = ln (0.693t1/2) D. E. k = 0.693t1/2
27. Uranium-235 has a half-life of 7.04 ´ 108 years. How many years will it take for 99.9% of a U-235 sample to decay? A. 7.0 ´ 105 yr B. 1.0 ´ 106 yr C. 4.7 ´ 109 yr D. 4.9 ´ 109 yr E. 7.0 ´ 109 yr
28. How long does it take for 105 mg of tritium, , to decay to 35.0 mg? The half-life of tritium is 12.3 yr. A. 12.3 yr B. 14.2 yr C. 19.5 yr D. 24.6 yr E. 36.9 yr
29. The half-life of mercury-203 is 46.6 days. What percentage of mercury-203 remains in a sample after 365 days? A. 3.72 ´ 10-4 % B. 0.439% C. 1.49% D. 5.43% E. 12.8%
30. How many atoms decay in 12 hours in a 3.5 ´ 10-8 Ci source? (1 Ci = 3.700 ´ 1010 atom/s) A. 4.1 ´ 10-14 atoms B. 1.6 ´ 104 atoms C. 5.0 ´ 104 atoms D. 5.6 ´ 107 atoms E. 4.6 ´ 1022 atoms
31. The half-life of americium-241 is 457 years. What is the activity in curies of a 0.010 g sample of Am-241? (1 Ci = 3.700 ´ 1010 atom/s) A. 0.032 Ci B. 1.0 Ci C. 1.0 ´ 106 Ci D. 1.0 ´ 109 Ci E. 3.79 ´ 1016 Ci
32. A 1.00 g sample of decays to 0.87 g in 1.9 hr. What is the half-life of ? A. 0.073 hr B. 0.11 hr C. 4.5 hr D. 9.5 hr E. 14 hr
33. decays by beta emission. It has a decay constant of 7.64 ´ 10-10 s-1. How many beta particles are emitted per second from 1.00 ´ 10-2 moles of ? A. 7.64 ´ 10-12 s-1 B. 7.64 ´ 10-10 s-1 C. 3.12 ´ 105 s-1 D. 1.45 ´ 1010 s-1 E. 4.60 ´ 1012 s-1
34. decays by alpha emission. The isotope's half-life is 2.4 ´ 104 yr. How many alpha particles are emitted per second by 1.0 ´ 10-4 mol . A. 9.2 ´ 10-13 s-1 B. 3.0 ´ 10-9 s-1 C. 5.5 ´ 107 s-1 D. 7.6 ´ 1011 s-1 E. 4.8 ´ 1016 s-1
35. The rate constant for the decay of copper-62 is 9.8 minutes-1. What is the half-life of this isotope? A. 0.071 m. B. 1.6 m. C. 14 m. D. 0.10 m. E. 9.8 m.
36. The half-life of is 9.5 ´ 1012 s. How many atoms of are present in a sample with an activity of 3.9 ´ 109 atoms/s? A. 1.9 ´ 10-23 atoms B. 4.1 ´ 10-4 atoms C. 2.4 ´ 103 atoms D. 9.1 ´ 1016 atoms E. 5.3 ´ 1022 atoms
37. The decay constant, k, for is 1.37 ´ 10-11 s-1. What mass of has an activity of 1.06 ´ 109 atoms/s? A. 6.60 ´ 10-6 g B. 1.49 ´ 10-3 g C. 2.90 ´ 10-2 g D. 3.44 ´ 101 g E. 8.97 ´ 102 g
38. is a naturally occurring radioisotope with a half-life of 7.1 ´ 108 yr. What percentage of that was present during the formation of Earth still remains? Assume that the Earth is 4.5 ´ 109 yr old and that no is created by decomposition of other elements. A. 0.16% B. 0.81% C. 1.2% D. 4.4% E. 6.3%
39. is used in many home smoke alarms. It has a decay half-life of 457 years. How long does it take for the americium in a smoke alarm to decay to 6.25% of its original mass? A. 870 yr B. 1300 yr C. 1426 yr D. 1828 yr E. 2600 yr
40. All of the following statements concerning radiocarbon dating are true EXCEPT A. radiocarbon dating is used to determine the age of meteorites. B. the decay of is used for radiocarbon dating. C. radioactive carbon is produced by the reaction of cosmic radiation and nitrogen-14. D. radioactive carbon decays into a nitrogen-14 atom and a beta particle. E. radiocarbon dating is limited to objects less than 50,000 years old.
41. A 1.00 g sample of carbon taken from a fossilized plant has an activity of 0.19 atoms/s. What mass of is present? The decay constant of is 3.8 ´ 10-12 s-1. A. 4.6 ´ 10-34 g B. 5.0 ´ 10-14 g C. 1.2 ´ 10-12 g D. 6.1 ´ 10-12 g E. 3.2 ´ 10-11 g
42. A 1.00 g sample of carbon taken from a living plant has an activity of 13.6 atoms/min. If 1.00 g of an organic sample is found to have an activity of 6.27 atoms/min, what is the age of the sample? The decay constant of is 1.21 ´ 10-4 yr-1. A. 2.64 ´ 103 yr B. 5.28 ´ 103 yr C. 6.40 ´ 103 yr D. 8.18 ´ 103 yr E. 1.07 ´ 104 yr
43. Carbon dating can be applied to organic materials up to 5.0 ´ 104 years old. What percentage of the original is present in a sample after this much time? The half-life of is 5730 yr. A. 0.0011% B. 0.24% C. 0.89% D. 6.0% E. 8.7%
44. An extremely sensitive method has recently been developed for measuring concentrations in organic materials. What is this method? A. acid titration B. nuclear magnetic resonance C. gas chromatography D. atomic absorption spectrometry E. mass spectrometry
45. The energy change for a nuclear reaction can be calculated from which of the following equations? A. B. DE = hv C. DE = cDm2 D. DE = c2Dm E.
46. The point of maximum stability in the binding energy curve occurs in the vicinity of which one of the following isotopes? A. B. C. D. E.
47. All of the following statements are true EXCEPT A. a nucleus weighs less than the protons and neutrons from which it is composed. B. the difference in mass between a nucleus and its individual protons and neutrons is called the mass defect. C. the process in which a nucleus decomposes to form two or more lighter nuclei is called nuclear fission. D. nuclear fission evolves more energy than ordinary chemical reactions, such as the combustion of gasoline. E. nuclear fission evolves more energy than nuclear fusion.
48. All of the following statements are true EXCEPT A. for a nuclear fission chain reaction to occur, one neutron must be produced for each one consumed. B. in a light water nuclear reactor, steam drives the turbines used to generate electricity. C. in a fission reactor, reacts with a neutron to produce more neutrons and products of lower mass. D. many of the fission products of are radioactive. E. one function of the water in a light water reactor is to slow down neutrons given off by fission.
49. The molar nuclear mass of fluorine-19 is 18.99840 g/mol. The molar mass of a proton is 1.007825 g/mol. The molar mass of a neutron is 1.008665 g/mol. Calculate the binding energy (in kJ/mol) of F-19. (c = 2.998 ´ 108 m/s) A. 6.753 ´ 109 kJ/mol B. 7.131 ´ 109 kJ/mol C. 1.426 ´ 1010 kJ/mol D. 8.609 ´ 1011 kJ/mol E. 8.538 ´ 1011 kJ/mol
50. Calculate the energy change per mole for the following reaction? The mass of is 2.01355 amu and the mass of is 4.00150 amu. A. 2.30 ´ 109 kJ/mol B. 9.00 ´ 109 kJ/mol C. 1.79 ´ 1011 kJ/mol D. 3.52 ´ 1012 kJ/mol E. 2.84 ´ 1013 kJ/mol
51. Calculate the energy released (per mole of deuterium consumed) for the following fusion reaction, given the following molar masses of nucleons and nuclei. (c = 2.998 ´ 108 m/s)
particle proton neutron deuterium tritium helium-4
mass (g/mol) 1.007825 1.008665 2.0140 3.01605 4.00260
A. 5.63 ´ 103 kJ/mol B. 1.69 ´ 1012 kJ/mol C. 4.62 ´ 1010 kJ/mol D. 8.44 ´ 108 kJ/mol E. 1.69 ´ 109 kJ/mol 52. The following fusion reaction releases 2.16 ´ 109 kJ per mole of deuterium reacted. How much mass is lost in the reaction per mole of deuterium reacted? A. 1.2 ´ 10-5 g B. 2.4 ´ 10-5 g C. 1.2 ´ 10-2 g D. 2.4 ´ 10-2 g E. 1.2 ´ 10-1 g
53. decays by positron emission. Calculate the energy released or gained in this reaction using the information below.
mass (amu) 0.00055 1.00867 1.00728 11.00656 11.00814
A. 1.03 ´ 10-3 kJ/mol evolved in reaction B. 9.27 ´ 107 kJ/mol evolved in reaction C. 1.42 ´ 108 kJ/mol evolved in reaction D. 9.12 ´ 109 kJ/mol absorbed in reaction E. 8.74 ´ 1013 kJ/mol absorbed in reaction 54. All of the following statements are true EXCEPT A. the Sun's energy comes primarily from the fusion of hydrogen. B. fusion reactions have high activation barriers due to repulsion of nuclei. C. fission reactions have relatively low activation barriers because uncharged neutrons used to initiate the reactions are not repelled by nuclei. D. magnetic fields can be used to confine nuclei during a fusion reaction. E. many fission reactors have been replaced by fusion reactors, which produce more energy and fewer waste products.
55. Human exposure to radiation comes primarily from A. medical x-rays. B. sitting too close to the television. C. nuclear fallout from bomb tests. D. radon gas. E. using cellular telephones.
56. Marie Curie was the first person in history to A. discover radioactivity. B. die from leukemia. C. win two Nobel Prizes. D. name an element. E. discover the neutron.
57. Most of the sun's energy is produced primarily from A. nuclear fission of 235U. B. nuclear fusion of 235U. C. nuclear fission of 1H. D. nuclear fusion of 1H. E. nuclear fusion of 2H.
58. Most of the nuclear power on earth is produced primarily from A. nuclear fission of 235U. B. nuclear fusion of 235U. C. nuclear fission of 1H. D. nuclear fusion of 1H. E. nuclear fusion of 2H.
59. Radioactive elements have been used A. to determine the age of artifacts. B. to treat food to reduce spoilage. C. in smoke detectors. D. both a and b E. all of the above
Chapter 19--Complex Ions 1. A central metal atom bonded to surrounding ions or molecules is called a A. coordination compound. B. ligand. C. chelate. D. Lewis acid. E. Lewis base.
2. During the formation of a coordination compound, the metal acts as a A. Brønsted-Lowry base. B. chelate. C. ligand. D. Lewis acid. E. Lewis base.
3. A chelating agent is a A. molecule or ion that contains nitrogen. B. ligand with a negative charge (e.g. CN-). C. ligand with no charge (e.g. H2O). D. another name for a central metal atom in a coordination compound. E. ligand that can form more than one bond with a central metal atom.
4. A Lewis base is defined as a species that A. increases the OH- concentration in water. B. accepts a proton. C. donates a pair of electrons to a bond. D. has a negative charge. E. is two electrons short of an octet in the valence shell.
5. Ions such as [Co(H2O)6]3+ and [Ag(CN)2]– are called ____. A. ligands B. Lewis bases C. chelates D. alloys E. coordination complexes
6. Which molecule or ion does NOT act as a ligand? A. H2O B. NH4+ C. Cl– D. C2O42– E. CN–
7. What is the oxidation number of the copper in [Cu(NH3)3Cl]NO3? A. 0 B. +1 C. +2 D. +4 E. +5
8. What is the oxidation number of the chromium in [Cr(H2O)2Cl4]-? A. 0 B. +1 C. +2 D. +3 E. +5
9. What is the oxidation number of molybdenum in [Mo(H2O)5OH]Cl2? A. +1 B. +2 C. +3 D. +4 E. +6
10. What is the oxidation number of the iron in [Fe(CN)6]3-? A. -3 B. -2 C. +2 D. +3 E. +4
11. What is the coordination number of the central metal ion in [Ni(NH3)4(H2O)2]Cl2? A. 0 B. 2 C. 4 D. 6 E. 8
12. What is the coordination number of the copper in [Cu(NH3)2]Cl? A. 1 B. 2 C. 3 D. 4 E. 8
13. What is the coordination number of the central metal ion in [Cr(en)2(OH)2]NO3? (en = H2NCH2CH2NH2) A. 0 B. 2 C. 4 D. 6 E. 7
14. Identify the ligands and their charges in [Os(NH3)4Cl2]+. A. ammonia (no charge), chloride ion (charge = -1) B. chloride ion (charge = -1) C. ammonia (no charge) D. osmium (charge = +3) E. osmium (charge = +3), chloride ion (charge = -1)
15. Identify the ligands and their charges in [Cr(en)4(CN)2]NO3. A. ethylenediamine (no charge) B. cyanide (charge = no charge) C. nitrate ion (charge = -1) D. ethylenediamine (no charge), cyanide ion (charge = -1) E. ethylenediamine (no charge), cyanide ion (charge = -1), nitrate ion (charge = -1)
16. Identify the ligands and their charges in [Co(en)2Br2]+. A. ethylene (charge = -1), bromide (charge = 0) B. ethylene (charge = -1), bromide ion (charge = -1) C. ethylenediamine (charge = 0), bromide ion (charge = -1), cobalt (charge = +3) D. ethylenediamine (charge = 0), bromide (charge = 0) E. ethylenediamine (charge = 0), bromide ion (charge = -1)
17. What is the mass percent of chromium in [Cr(C2O4)2]SO4? A. 16.0% B. 21.3% C. 22.8% D. 36.5% E. 67.5%
18. What is the mass percent of chromium in the sulfate salt of [Cr(H2O)5(OH)]2+? A. 20.38% B. 27.20% C. 32.68% D. 67.32% E. 72.80%
19. What is the correct formula for a nickel(II) complex which contains 4 ammonia molecules and 2 water molecules? A. Ni(NH3)4(H2O)2 B. Ni(NH3)4(H2O)22+ C. [Ni(NH3)4](H2O)2 D. Ni(NH3)4(H2O)22E. Ni(NH3)4(H2O)24-
20. What is the correct formula for a square planar copper(II) complex that is bonded to ethylenediamine ligands? A. Cu(en)2+ B. Cu(en)22+ C. Cu(en)32+ D. Cu(en)42+ E. Cu(en)62+
21. What is the name of the compound having the formula K2[PtCl4]? A. potassium chloroplatinate(II) B. potassium tetrachloroplatinate(II) C. potassium chloroplatinate(IV) D. potassium platanotetrachlorate(II) E. dipotassium tetrachloroplatnum(II)
22. What is the name of the compound having the formula [Cr(en)2(H2O)2]SO4? A. dihydroxydiethlyenediamminechromate(II) sulfate B. diaquabis(ethylenediamine)sulfatochromate(IV) C. bis(ethylenediamine)diaquachromium(II) sulfato D. diaquabis(ethylenediamine)sulfatochromium(II) E. diaquabis(ethylenediamine)chromium(II) sulfate
23. What is the name of the compound having the formula (NH4)2[Cu(CN)4]? A. ammonium tetracyanocuprate(II) B. diammonium tetracyanocuprate(II) C. bisamminetetracyanocuprate(II) D. tetracyanobisamminecuprate(II) E. tetracyanocuprate(II) ammonia
24. What is the formula for dicyanobis(ethylenediamine)zirconium(IV) nitrate? A. Zr[(CN)2(NO3)2(en)2] B. [Zr(CN)2(NO3)2(en)2] C. (NO3)2[Zr(CN)2(en)2] D. [Zr(CN)2(en)2](NO3)2 E. [Zr(NO3)2(en)2](CN)2
25. What is the formula for potassium diamminetetrachlorovanadate(III)? A. [KV(NH3)2Cl4] B. K2[V(NH3)2Cl4] C. K[V(NH3)2Cl4] D. [KV(NH3)2]Cl4 E. K[V(NH3)2]Cl4
26. Chlorophylls are coordination compounds that contain a magnesium ion as the central metal. The most abundant chlorophyll, called chlorophyll a, has a mass percent of magnesium of 2.72%. What is the molar mass of chlorophyll a? A. 2.72 g/mol B. 57.2 g/mol C. 112 g/mol D. 583 g/mol E. 893 g/mol
27. Which of the following molecules or ions are chelating agents: NH3, H2NCH2CH2NH2, CO2, and C2O42-? A. NH3 B. NH3 and CO2 C. CO2 and H2NCH2CH2NH2 D. H2NCH2CH2NH2 and C2O42E. CO2, H2NCH2CH2NH2 and C2O42-
28. What are the possible formulas of all the tetracoordinate complex ions formed by copper(II) with NH3 and/or ethylenediamine (en) molecules as ligands? A. Cu(NH3)42+, Cu(en)(NH3)32+, Cu(en)2(NH3)22+, Cu(en)3(NH3)2+, Cu(en)42+ B. Cu(NH3)42+, Cu(en)2(NH3)42+, Cu(en)22+ C. Cu(NH3)42+, Cu(en)(NH3)22+, Cu(en)22+ D. Cu(NH3)62+, Cu(en)2(NH3)22+, Cu(en)32+ E. Cu(NH3)42+, Cu(en)22+
29. What are the possible geometries of a metal complex with a coordination number of 4? 1. 2. 3.
square planar tetrahedral octahedral
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 30. How many possible structures exist for the tetrahedral complex [Zn(en)3]2+? A. 1 B. 2 C. 3 D. 4 E. 6
31. Which of the following species have geometric isomers: [Fe(en)Cl4]–, [Fe(en)2Cl2]+, and [Fe(en)2BrCl]+? A. [Fe(en)Cl4]– only B. [Fe(en)2Cl2]+ only C. [Fe(en)2BrCl]+ only D. [Fe(en)2Cl2]+ and [Fe(en)2BrCl]+ E. [Fe(en)Cl4]–, [Fe(en)2Cl2]+, and [Fe(en)2BrCl]+
32. How many geometric isomers may exist for the square-planar complex ion [Pt(OH)2(NH3)2]2–? A. 2 B. 3 C. 4 D. 5 E. 6
33. How many geometric isomers exist for [Fe(en)2(CN)2]+? A. 1 B. 2 C. 3 D. 4 E. 6
34. How many geometric isomers exist for [Fe(H2O)4(NH3)2]3+? A. 1 B. 2 C. 3 D. 4 E. 6
35. Potassium hexacyanoiron(II), K4[Fe(CN)6], is water soluble. In an aqueous solution, it will decompose into A. 2 ions. B. 4 ions. C. 5 ions. D. 6 ions. E. 10 ions.
36. What is the electron configuration of Cr2+? A. 1s22s22p63s23p6 B. 1s22s22p63s23p63d64s2 C. 1s22s22p63s23p63d34s2 D. 1s22s22p63s23p63d54s1 E. 1s22s22p63s23p63d4
37. What is the electron configuration for Ti4+? A. 1s22s22p63s23p64s2 B. 1s22s22p63s23p63d2 C. 1s22s22p63s23p63d24s2 D. 1s22s22p63s23p63d10 E. 1s22s22p63s23p6
38. What is the electron configuration for Fe3+? A. 1s22s22p63s23p63d54s1 B. 1s22s22p63s23p63d44s2 C. 1s22s22p63s23p63d6 D. 1s22s22p63s23p63d5 E. 1s22s22p63s23p63d7
39. What is the correct orbital diagram for Mn2+?
1) 2) 3) 4) 5)
A. 1 B. 2 C. 3 D. 4 E. 5
1s (¯) (¯) (¯) (¯) (¯)
2s (¯) (¯) (¯) (¯) (¯)
2p (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯)
3s (¯) (¯) (¯) (¯) (¯)
3p (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯)
3d
4s
( )( )( )( )( ) ( )( )( )( )( ) ( )( )( )( )( ) ( )( )( )( )( )
( ) (¯) (¯)
40. What is the correct orbital diagram for Cr2+
1) 2) 3) 4) 5)
1s (¯) (¯) (¯) (¯) (¯)
2s (¯) (¯) (¯) (¯) (¯)
2p (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯)
3s (¯) (¯) (¯) (¯) (¯)
3p (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯)
3d ( )( )( )( )( ) ( )( )( )( )( ) ( )( )( )( )( ) ( )( )( )( )( ) ( )( )( )( )( )
4s
(¯) (¯)
A. 1 B. 2 C. 3 D. 4 E. 5 41. All of the following statements concerning crystal field theory are true EXCEPT A. in low-spin complexes, electrons are concentrated in the dxy, dyz, and dxz orbitals. B. in an isolated atom or ion, the five d orbitals have identical energy. C. low-spin complexes contain the maximum number of unpaired electrons. D. the crystal field splitting is larger in low-spin complexes than high-spin complexes. E. the energy difference between d orbitals often corresponds to an energy of visible light.
42. When 6 ligands approach a central metal atom along the x-, y- and z-axes, they will cause the d orbitals to split in energy. The and the orbitals are most affected because A. the electron density in these orbitals is squared. B. unlike the other d orbitals, these orbitals contain unpaired electrons. C. unlike the other d orbitals, these orbitals each contain a pair of electrons. D. they have the highest electron density along the x-, y-, and z-axes, and thus feel the greatest repulsion from the ligands' unpaired electrons. E. the shape of these orbitals is similar to that of the ligands' orbitals.
43. All of the followings statements are true EXCEPT A. transition metal complexes are often colored, due to 3d to 1s electronic transitions. B. in low-spin complexes, electrons are concentrated in the lower-energy orbitals. C. low-spin complexes contain the minimum number of unpaired electrons. D. the crystal field splitting energy is small in high-spin complexes. E. the crystal field splitting energy of a metal complex can sometimes be determined from the absorption spectrum.
44. The complex FeF64- is paramagnetic. Which of the following set of terms best describes the complex? A. high-spin, Do large B. high-spin, Do small C. low-spin, Do large D. low-spin, Do small E. none of the above
45. Determine the number of unpaired electrons in an octahedral, high-spin Cr(II) complex. A. 1 B. 2 C. 3 D. 4 E. 5
46. Determine the number of unpaired electrons in the low-spin complex ion Cr(CN)64-. A. 0 B. 1 C. 2 D. 4 E. 5
47. Determine the number of unpaired electrons in an octahedral, strong-field Mn2+ complex ion. A. 1 B. 2 C. 3 D. 4 E. 5
48. If an octahedral iron(II) complex is diamagnetic, which of the following sets of conditions best describes the complex? A. low-spin, D0 small B. low-spin, D0 large C. high-spin, D0 small D. high-spin, D0 large E. none of these
49. Which of the following ligands causes the greatest splitting of d orbitals? A. H2O B. NO2C. ClD. FE. NH3
50. Which of the following ligands causes the smallest splitting of d orbitals? A. ClB. en C. CND. NH3 E. NO3-
51. All of the following complexes will be colored (in the visible) EXCEPT ____. A. [Co(en)3]2+ B. [Fe(CN)6]4– C. [Ni(H2O)6]2+ D. [Sc(NH3)6]3+ E. [Cu(NH3)4]2+
52. All of the following ions are able to form both low-spin and high-spin octahedral complex ions EXCEPT A. Cr2+. B. Cr3+. C. Mn2+. D. Fe3+. E. Co3+.
53. The spectrochemical series A. lists metals ions in order of their oxidation state. B. lists metals ions in order by their color. C. lists ligands in order of their tendency to split d orbitals. D. is used to determine the oxidation state of a central metal atom. E. is used to calculate the crystal field splitting energy of complex ions.
54. Why might a cyano complex of a metal appear yellow and the chloride complex of the same metal appear green? A. The cyanide ions transmits yellow light more efficiently than the chloride ions. B. The cyanide ions splits the d orbital energies more than the chloride ions, resulting in the cyanide complex absorbing yellow light. C. The cyanide ions splits the d orbital energies more than chloride ions, resulting in the absorption band for the cyanide complex being shifted to shorter wavelengths. D. The chloride ions splits the d orbital energies more than cyanide ions, resulting in the absorption band for the chloride complex being shifted to shorter wavelengths. E. The chloride ions splits the d orbital energies more than the cyanide ions, resulting in the chloride complex absorbing yellow light.
55. The wavelength of maximum absorption of Cu(NH3)42+ is 5.80 ´ 102 nm. What is the energy of the electronic transition? A. 3.80 ´ 10-36 J B. 1.16 ´ 10-31 J C. 1.16 ´ 10-22 J D. 3.42 ´ 10-19 J E. 2.63 ´ 10+35 J
56. The crystal field splitting energy for TiF63- is 203 kJ/mol. At what wavelength does the maximum absorbance of light occur for this complex? A. 9.80 ´ 10-31 nm B. 9.80 ´ 10-28 nm C. 4.32 ´ 102 nm D. 5.90 ´ 102 nm E. 1.02 ´ 103 nm
57. Ti(NH3)63+ has a d-orbital electron transition at 399 nm. Find Do At this wavelength. A. 300 kJ/mol B. 235 kJ/mol C. 499 kJ/mol D. 399 kJ/mol E. 150 kJ/mol
58. Alfred Werner published his theory of coordination chemistry in 1893. One of the postulates of his theory was that coordination complexes have "primary valences" and "secondary valences." He showed this by testing the ____ of solutions containing [Co(NH3)6]Cl3 vs. [Co(NH3)5 Cl]Cl2 vs. [Co(NH3)4 Cl2]Cl. A. color B. conductivity C. density D. melting point E. taste
59. Heme is a coordination compound essential to your life. What is the metal ion coordinated in the heme of hemoglobin? A. Cu2+ B. Ca2+ C. Co2+ D. Fe2+ E. Mg2+
Chapter 19--Complex Ions Key
1. A central metal atom bonded to surrounding ions or molecules is called a A. coordination compound. B. ligand. C. chelate. D. Lewis acid. E. Lewis base.
2. During the formation of a coordination compound, the metal acts as a A. Brønsted-Lowry base. B. chelate. C. ligand. D. Lewis acid. E. Lewis base.
3. A chelating agent is a A. molecule or ion that contains nitrogen. B. ligand with a negative charge (e.g. CN-). C. ligand with no charge (e.g. H2O). D. another name for a central metal atom in a coordination compound. E. ligand that can form more than one bond with a central metal atom.
4. A Lewis base is defined as a species that A. increases the OH- concentration in water. B. accepts a proton. C. donates a pair of electrons to a bond. D. has a negative charge. E. is two electrons short of an octet in the valence shell.
5. Ions such as [Co(H2O)6]3+ and [Ag(CN)2]– are called ____. A. ligands B. Lewis bases C. chelates D. alloys E. coordination complexes
6. Which molecule or ion does NOT act as a ligand? A. H2O B. NH4+ C. Cl– D. C2O42– E. CN–
7. What is the oxidation number of the copper in [Cu(NH3)3Cl]NO3? A. 0 B. +1 C. +2 D. +4 E. +5
8. What is the oxidation number of the chromium in [Cr(H2O)2Cl4]-? A. 0 B. +1 C. +2 D. +3 E. +5
9. What is the oxidation number of molybdenum in [Mo(H2O)5OH]Cl2? A. +1 B. +2 C. +3 D. +4 E. +6
10. What is the oxidation number of the iron in [Fe(CN)6]3-? A. -3 B. -2 C. +2 D. +3 E. +4
11. What is the coordination number of the central metal ion in [Ni(NH3)4(H2O)2]Cl2? A. 0 B. 2 C. 4 D. 6 E. 8
12. What is the coordination number of the copper in [Cu(NH3)2]Cl? A. 1 B. 2 C. 3 D. 4 E. 8
13. What is the coordination number of the central metal ion in [Cr(en)2(OH)2]NO3? (en = H2NCH2CH2NH2) A. 0 B. 2 C. 4 D. 6 E. 7
14. Identify the ligands and their charges in [Os(NH3)4Cl2]+. A. ammonia (no charge), chloride ion (charge = -1) B. chloride ion (charge = -1) C. ammonia (no charge) D. osmium (charge = +3) E. osmium (charge = +3), chloride ion (charge = -1)
15. Identify the ligands and their charges in [Cr(en)4(CN)2]NO3. A. ethylenediamine (no charge) B. cyanide (charge = no charge) C. nitrate ion (charge = -1) D. ethylenediamine (no charge), cyanide ion (charge = -1) E. ethylenediamine (no charge), cyanide ion (charge = -1), nitrate ion (charge = -1)
16. Identify the ligands and their charges in [Co(en)2Br2]+. A. ethylene (charge = -1), bromide (charge = 0) B. ethylene (charge = -1), bromide ion (charge = -1) C. ethylenediamine (charge = 0), bromide ion (charge = -1), cobalt (charge = +3) D. ethylenediamine (charge = 0), bromide (charge = 0) E. ethylenediamine (charge = 0), bromide ion (charge = -1)
17. What is the mass percent of chromium in [Cr(C2O4)2]SO4? A. 16.0% B. 21.3% C. 22.8% D. 36.5% E. 67.5%
18. What is the mass percent of chromium in the sulfate salt of [Cr(H2O)5(OH)]2+? A. 20.38% B. 27.20% C. 32.68% D. 67.32% E. 72.80%
19. What is the correct formula for a nickel(II) complex which contains 4 ammonia molecules and 2 water molecules? A. Ni(NH3)4(H2O)2 B. Ni(NH3)4(H2O)22+ C. [Ni(NH3)4](H2O)2 D. Ni(NH3)4(H2O)22E. Ni(NH3)4(H2O)24-
20. What is the correct formula for a square planar copper(II) complex that is bonded to ethylenediamine ligands? A. Cu(en)2+ B. Cu(en)22+ C. Cu(en)32+ D. Cu(en)42+ E. Cu(en)62+
21. What is the name of the compound having the formula K2[PtCl4]? A. potassium chloroplatinate(II) B. potassium tetrachloroplatinate(II) C. potassium chloroplatinate(IV) D. potassium platanotetrachlorate(II) E. dipotassium tetrachloroplatnum(II)
22. What is the name of the compound having the formula [Cr(en)2(H2O)2]SO4? A. dihydroxydiethlyenediamminechromate(II) sulfate B. diaquabis(ethylenediamine)sulfatochromate(IV) C. bis(ethylenediamine)diaquachromium(II) sulfato D. diaquabis(ethylenediamine)sulfatochromium(II) E. diaquabis(ethylenediamine)chromium(II) sulfate
23. What is the name of the compound having the formula (NH4)2[Cu(CN)4]? A. ammonium tetracyanocuprate(II) B. diammonium tetracyanocuprate(II) C. bisamminetetracyanocuprate(II) D. tetracyanobisamminecuprate(II) E. tetracyanocuprate(II) ammonia
24. What is the formula for dicyanobis(ethylenediamine)zirconium(IV) nitrate? A. Zr[(CN)2(NO3)2(en)2] B. [Zr(CN)2(NO3)2(en)2] C. (NO3)2[Zr(CN)2(en)2] D. [Zr(CN)2(en)2](NO3)2 E. [Zr(NO3)2(en)2](CN)2
25. What is the formula for potassium diamminetetrachlorovanadate(III)? A. [KV(NH3)2Cl4] B. K2[V(NH3)2Cl4] C. K[V(NH3)2Cl4] D. [KV(NH3)2]Cl4 E. K[V(NH3)2]Cl4
26. Chlorophylls are coordination compounds that contain a magnesium ion as the central metal. The most abundant chlorophyll, called chlorophyll a, has a mass percent of magnesium of 2.72%. What is the molar mass of chlorophyll a? A. 2.72 g/mol B. 57.2 g/mol C. 112 g/mol D. 583 g/mol E. 893 g/mol
27. Which of the following molecules or ions are chelating agents: NH3, H2NCH2CH2NH2, CO2, and C2O42-? A. NH3 B. NH3 and CO2 C. CO2 and H2NCH2CH2NH2 D. H2NCH2CH2NH2 and C2O42E. CO2, H2NCH2CH2NH2 and C2O42-
28. What are the possible formulas of all the tetracoordinate complex ions formed by copper(II) with NH3 and/or ethylenediamine (en) molecules as ligands? A. Cu(NH3)42+, Cu(en)(NH3)32+, Cu(en)2(NH3)22+, Cu(en)3(NH3)2+, Cu(en)42+ B. Cu(NH3)42+, Cu(en)2(NH3)42+, Cu(en)22+ C. Cu(NH3)42+, Cu(en)(NH3)22+, Cu(en)22+ D. Cu(NH3)62+, Cu(en)2(NH3)22+, Cu(en)32+ E. Cu(NH3)42+, Cu(en)22+
29. What are the possible geometries of a metal complex with a coordination number of 4? 1. 2. 3.
square planar tetrahedral octahedral
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 30. How many possible structures exist for the tetrahedral complex [Zn(en)3]2+? A. 1 B. 2 C. 3 D. 4 E. 6
31. Which of the following species have geometric isomers: [Fe(en)Cl4]–, [Fe(en)2Cl2]+, and [Fe(en)2BrCl]+? A. [Fe(en)Cl4]– only B. [Fe(en)2Cl2]+ only C. [Fe(en)2BrCl]+ only D. [Fe(en)2Cl2]+ and [Fe(en)2BrCl]+ E. [Fe(en)Cl4]–, [Fe(en)2Cl2]+, and [Fe(en)2BrCl]+
32. How many geometric isomers may exist for the square-planar complex ion [Pt(OH)2(NH3)2]2–? A. 2 B. 3 C. 4 D. 5 E. 6
33. How many geometric isomers exist for [Fe(en)2(CN)2]+? A. 1 B. 2 C. 3 D. 4 E. 6
34. How many geometric isomers exist for [Fe(H2O)4(NH3)2]3+? A. 1 B. 2 C. 3 D. 4 E. 6
35. Potassium hexacyanoiron(II), K4[Fe(CN)6], is water soluble. In an aqueous solution, it will decompose into A. 2 ions. B. 4 ions. C. 5 ions. D. 6 ions. E. 10 ions.
36. What is the electron configuration of Cr2+? A. 1s22s22p63s23p6 B. 1s22s22p63s23p63d64s2 C. 1s22s22p63s23p63d34s2 D. 1s22s22p63s23p63d54s1 E. 1s22s22p63s23p63d4
37. What is the electron configuration for Ti4+? A. 1s22s22p63s23p64s2 B. 1s22s22p63s23p63d2 C. 1s22s22p63s23p63d24s2 D. 1s22s22p63s23p63d10 E. 1s22s22p63s23p6
38. What is the electron configuration for Fe3+? A. 1s22s22p63s23p63d54s1 B. 1s22s22p63s23p63d44s2 C. 1s22s22p63s23p63d6 D. 1s22s22p63s23p63d5 E. 1s22s22p63s23p63d7
39. What is the correct orbital diagram for Mn2+?
1) 2) 3) 4) 5)
A. 1 B. 2 C. 3 D. 4 E. 5
1s (¯) (¯) (¯) (¯) (¯)
2s (¯) (¯) (¯) (¯) (¯)
2p (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯)
3s (¯) (¯) (¯) (¯) (¯)
3p (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯)
3d
4s
( )( )( )( )( ) ( )( )( )( )( ) ( )( )( )( )( ) ( )( )( )( )( )
( ) (¯) (¯)
40. What is the correct orbital diagram for Cr2+
1) 2) 3) 4) 5)
1s (¯) (¯) (¯) (¯) (¯)
2s (¯) (¯) (¯) (¯) (¯)
2p (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯)
3s (¯) (¯) (¯) (¯) (¯)
3p (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯) (¯)(¯)(¯)
3d ( )( )( )( )( ) ( )( )( )( )( ) ( )( )( )( )( ) ( )( )( )( )( ) ( )( )( )( )( )
4s
(¯) (¯)
A. 1 B. 2 C. 3 D. 4 E. 5 41. All of the following statements concerning crystal field theory are true EXCEPT A. in low-spin complexes, electrons are concentrated in the dxy, dyz, and dxz orbitals. B. in an isolated atom or ion, the five d orbitals have identical energy. C. low-spin complexes contain the maximum number of unpaired electrons. D. the crystal field splitting is larger in low-spin complexes than high-spin complexes. E. the energy difference between d orbitals often corresponds to an energy of visible light.
42. When 6 ligands approach a central metal atom along the x-, y- and z-axes, they will cause the d orbitals to split in energy. The and the orbitals are most affected because A. the electron density in these orbitals is squared. B. unlike the other d orbitals, these orbitals contain unpaired electrons. C. unlike the other d orbitals, these orbitals each contain a pair of electrons. D. they have the highest electron density along the x-, y-, and z-axes, and thus feel the greatest repulsion from the ligands' unpaired electrons. E. the shape of these orbitals is similar to that of the ligands' orbitals.
43. All of the followings statements are true EXCEPT A. transition metal complexes are often colored, due to 3d to 1s electronic transitions. B. in low-spin complexes, electrons are concentrated in the lower-energy orbitals. C. low-spin complexes contain the minimum number of unpaired electrons. D. the crystal field splitting energy is small in high-spin complexes. E. the crystal field splitting energy of a metal complex can sometimes be determined from the absorption spectrum.
44. The complex FeF64- is paramagnetic. Which of the following set of terms best describes the complex? A. high-spin, Do large B. high-spin, Do small C. low-spin, Do large D. low-spin, Do small E. none of the above
45. Determine the number of unpaired electrons in an octahedral, high-spin Cr(II) complex. A. 1 B. 2 C. 3 D. 4 E. 5
46. Determine the number of unpaired electrons in the low-spin complex ion Cr(CN)64-. A. 0 B. 1 C. 2 D. 4 E. 5
47. Determine the number of unpaired electrons in an octahedral, strong-field Mn2+ complex ion. A. 1 B. 2 C. 3 D. 4 E. 5
48. If an octahedral iron(II) complex is diamagnetic, which of the following sets of conditions best describes the complex? A. low-spin, D0 small B. low-spin, D0 large C. high-spin, D0 small D. high-spin, D0 large E. none of these
49. Which of the following ligands causes the greatest splitting of d orbitals? A. H2O B. NO2C. ClD. FE. NH3
50. Which of the following ligands causes the smallest splitting of d orbitals? A. ClB. en C. CND. NH3 E. NO3-
51. All of the following complexes will be colored (in the visible) EXCEPT ____. A. [Co(en)3]2+ B. [Fe(CN)6]4– C. [Ni(H2O)6]2+ D. [Sc(NH3)6]3+ E. [Cu(NH3)4]2+
52. All of the following ions are able to form both low-spin and high-spin octahedral complex ions EXCEPT A. Cr2+. B. Cr3+. C. Mn2+. D. Fe3+. E. Co3+.
53. The spectrochemical series A. lists metals ions in order of their oxidation state. B. lists metals ions in order by their color. C. lists ligands in order of their tendency to split d orbitals. D. is used to determine the oxidation state of a central metal atom. E. is used to calculate the crystal field splitting energy of complex ions.
54. Why might a cyano complex of a metal appear yellow and the chloride complex of the same metal appear green? A. The cyanide ions transmits yellow light more efficiently than the chloride ions. B. The cyanide ions splits the d orbital energies more than the chloride ions, resulting in the cyanide complex absorbing yellow light. C. The cyanide ions splits the d orbital energies more than chloride ions, resulting in the absorption band for the cyanide complex being shifted to shorter wavelengths. D. The chloride ions splits the d orbital energies more than cyanide ions, resulting in the absorption band for the chloride complex being shifted to shorter wavelengths. E. The chloride ions splits the d orbital energies more than the cyanide ions, resulting in the chloride complex absorbing yellow light.
55. The wavelength of maximum absorption of Cu(NH3)42+ is 5.80 ´ 102 nm. What is the energy of the electronic transition? A. 3.80 ´ 10-36 J B. 1.16 ´ 10-31 J C. 1.16 ´ 10-22 J D. 3.42 ´ 10-19 J E. 2.63 ´ 10+35 J
56. The crystal field splitting energy for TiF63- is 203 kJ/mol. At what wavelength does the maximum absorbance of light occur for this complex? A. 9.80 ´ 10-31 nm B. 9.80 ´ 10-28 nm C. 4.32 ´ 102 nm D. 5.90 ´ 102 nm E. 1.02 ´ 103 nm
57. Ti(NH3)63+ has a d-orbital electron transition at 399 nm. Find Do At this wavelength. A. 300 kJ/mol B. 235 kJ/mol C. 499 kJ/mol D. 399 kJ/mol E. 150 kJ/mol
58. Alfred Werner published his theory of coordination chemistry in 1893. One of the postulates of his theory was that coordination complexes have "primary valences" and "secondary valences." He showed this by testing the ____ of solutions containing [Co(NH3)6]Cl3 vs. [Co(NH3)5 Cl]Cl2 vs. [Co(NH3)4 Cl2]Cl. A. color B. conductivity C. density D. melting point E. taste
59. Heme is a coordination compound essential to your life. What is the metal ion coordinated in the heme of hemoglobin? A. Cu2+ B. Ca2+ C. Co2+ D. Fe2+ E. Mg2+
Chapter 20--Chemistry of the Metals 1. An ore is A. a natural source from which a metal can be extracted. B. a soluble mixture of ionic compounds. C. a mixture of two or more metallic elements. D. another name for a metal alloy, such as brass. E. a molten mixture of ionic compounds.
2. The principal ores of group 2 metals are A. carbonates. B. chlorides. C. oxides. D. phosphates. E. sulfides.
3. The principal ores of copper, silver and mercury are A. carbonates. B. chlorides. C. oxides. D. phosphates. E. sulfides.
4. Aluminum is produced by electrolysis of a mixture of bauxite ore (Al2O3) and cryolite (Na3AlF6). The purpose of the cryolite is to A. increase the current. B. decrease the potential at which aluminum is reduced. C. convert oxide ion into water D. lower the melting point. E. decrease the pH of the molten solution.
5. The main group metals in Group 1 are referred to as the A. alkali metals. B. alkaline earth metals. C. highly reactive metals. D. transition metals. E. basic metals.
6. The main group metals in Group 2 are referred to as the A. alkali metals. B. alkaline earth metals. C. less reactive metals. D. divalent metals. E. basic metals.
7. Write a balanced equation for the reduction of Co2O3 by carbon monoxide. A. Co2O3(s) + CO(g) ® 2Co(s) + CO4(g) B. 2Co2O3(s) + 2CO(g) ® 2Co2CO3(s) +O2(g) C. Co2O3(s) + CO(g) ® Co(s) + CO2(g) D. Co2O3(s) + CO(g) ® Co(s) + C(g) + 2O2(g) E. Co2O3(s) + 3CO(g) ® 2Co(s) + 3CO2(g)
8. Write a balanced equation for the reduction of manganese(IV) oxide ore by carbon monoxide. A. MnO(s) + CO(g) ® Mn(s) + CO2(g) B. MnO(s) + 2CO(g) ® Mn(s) + 2CO2(g) C. Mn2O(s) + CO(g) ® Mn(s) + CO2(g) D. MnO2(s) + 2CO(g) ® Mn(s) + 2CO2(g) E. MnO2(s) + 2CO(g) ® Mn(s) + 2C(s) + O2(g)
9. Both sodium metal and chlorine gas are created by the electrolysis of molten sodium chloride. If 1.00 ´ 102 amperes of current is passed through the cathode for 8.00 hr, what mass of Cl2 will be formed? (F = 96480 C/mol e-). A. 0.588 g B. 1.18 g C. 17.6 g D. 1.06 ´ 103 g E. 2.12 ´ 102 g
10. What is a useful by-product of the production of sodium metal? A. hydrogen B. magnesium C. chlorine D. helium E. calcium
11. Who worked out the process for obtaining aluminum from bauxite ore? A. Michael Faraday B. Enrico Fermi C. Charles Hall D. Linus Pauling E. Alfred Werner
12. Aluminum is obtained by A. the electrolysis of cryolite, Na3AlF6. B. reduction by a strong inorganic reducing agent, such as thiosulfate ion, S2O32-. C. the reduction of Al2O3 by carbon monoxide in a blast furnace. D. chemical decomposition of Al2O3 at temperatures greater than 2000C. E. electrolysis of a molten mixture of bauxite ore and cryolite, Na3AlF6.
13. What is the most common chemical reducing agent in metallurgical processes? A. thiosulfate ion, S2O32B. oxygen, O2 C. sodium metal D. carbon; in the form of carbon monoxide E. carbon; in the form of carbon dioxide
14. Write a balanced chemical equation for the chemical reduction of iron(III) oxide that occurs in a blast furnace. A. Fe2O3(s) + 3CO(g) ® 2Fe(s) + 3CO2(g) B. Fe2O3(s) + 3CO2(g) ® 2Fe(s) + 3CO(g) C. Fe2O3(s) + 3C(g) ® 2Fe(s) + 3CO(g) D. FeO(s) + CO2(g) ® FeCO3(s) E. FeO(s) + CO(g) ® Fe(s) + CO2(g)
15. At temperatures greater than 800C, limestone decomposes according to which balanced chemical equation? A. CaO(s) ® Ca(s) + O(g) B. 2CaO(s) ® 2Ca(s) + O2(g) C. CaCO3(s) ® CaO(s) + CO2(g) D. CaCO3(s) ® CaO2(s) + CO(g) E. CaCO2(g) ® Ca(s) + CO2(g)
16. Slag is a glassy material that forms when calcium oxide reacts with impurities (mostly silicon dioxide) in a blast furnace. Write the balanced chemical equation for the reaction. A. CaCO3(s) + Si(s) ® CaSiO3(l) B. CaCO3(s) + SiO2(s) ® CaSiO5(l) C. CaO(s) + SiO2(s) ® CaSiO(l) + O2(g) D. CaO(s) + SiO2(s) ® CaSiO3(l) E. CaO(s) + SiO(s) ® CaCO2(l)
17. Iron that is made in a blast furnace is called pig iron. The pig iron is converted into steel by reducing its carbon content from roughly 4% to under 2% by A. reduction of the carbon with hydrogen gas. B. the addition of calcium oxide to form slag. C. vacuum filtration. D. electrolysis. E. reaction with oxygen to form CO2(g).
18. The iron produced in a blast furnace contains significant impurities of carbon, silicon, manganese, and phosphorus. These impurities are removed in a ________. A. basic oxygen furnace B. distillation column C. mass spectrometer D. electrolytic cell E. centrifuge
19. Copper can be obtained from copper sulfide by heating in the presence of oxygen. This process is called A. roasting. B. oxidation. C. smelting. D. steaming. E. alloying.
20. The roasting of Cu2S(s) produces Cu(s) and SO2(g). What mass of sulfur dioxide is produced by the roasting of 1.00 ´ 103 kg Cu2S(s)? Cu2S(s) + O2(g) ® 2Cu(s) + SO2(g) A. 201 B. 402 kg C. 798 kg D. 804 kg E. 2.48 ´ 103 kg
21. Write a balanced chemical equation for the reaction of HgS with O2. A. HgS(s) + O2(g) ® HgO2(s) + S(s) B. HgS(s) + O2(g) ® Hg(l) + SO2(g) C. HgS(s) + O2(g) ® HgSO2(s) D. HgS(s) + 2O2(g) ® HgSO4(s) E. 2HgS(s) + O2(g) ® HgO(s) + 2S(s)
22. All of the following metals may be found in nature as free elements EXCEPT ____. A. Ir B. Au C. Pt D. Ti E. Rh
23. Separation of gold from ore is currently accomplished by using which ligand to make a soluble coordination compound from Au+1? A. chloride ion B. cyanide ion C. ammonia D. perchlorate ion E. ethylenediamine
24. Historically, miners separated gold by taking advantage of its A. high density. B. low melting point. C. high solubility in water. D. radioactivity. E. amphoteric nature.
25. Write a balanced equation for the reaction of sodium and bromine. A. Na(s) + Br-(l) ® NaBr(s) B. Na(s) + Br(l) ® NaBr(s) + Br-(aq) C. Na(s) + Br2(l) ® NaBr2(s) D. 2Na(s) + Br2(l) ® 2NaBr(s) E. 4Na(s)) + Br2(l) ® 2Na2Br(s)
26. Write a balanced chemical equation for the reaction of potassium and water. A. K(s) + H2O(l) ® KO(s) + H2(g) B. 2K(s) + 2H2O(l) ® 2KOH(aq) + H2(g) C. 2K(s) + 2H2O(l) ® K2O2(s) + 2H2(g) D. 2K(s) + H2O(l) ® K2O(s) + H2(g) E. 4K(s) + 2H2O(l) ® 4KH(s) + O2(g)
27. Write a balanced equation for the reaction of barium and nitrogen gas. A. Ba(s) + N2(g) ® BaN2(s) B. 2Ba(s) + N2(g) ® 2BaN(s) C. 2Ba(s) + N2(g) ® 2BaN(s) D. 3Ba(s) + N2(g) ® Ba3N2(s) E. 4Ba(s) + 3N2(g) ® 2Ba2N3(s)
28. Which of the following metals will react most violently with water? A. Li B. Mg C. K D. Ca E. Cr
29. Write a balanced equation for the reaction of Na and O2 to form sodium peroxide. A. 4Na(s) + O2(g) ® 2Na2O(s) B. Na(s) + O2(g) ® NaO2(s) C. 2Na(s) + O2(g) ® Na2O2(s) D. 2Na(s) + O2(g) ® 2NaO(s) E. 4Na(s) + 3O2(g) ® 2Na2O3(s)
30. The principal product of the reaction of potassium and oxygen is not K2O. The product is named ____ and has the formula ____. A. dipotassium trioxide, K2O3 B. potassium monoxide, KO C. potassium superoxide, KO2 D. potassium peroxide, K2O2 E. potassium trioxide, KO3
31. Write a balanced equation for the reaction of magnesium and nitrogen that occurs upon heating. A. Mg(s) + N2(g) ® MgN2(s) B. 2Mg(s) + N2(g) ® 2MgN(s) C. 4Mg(s) + N2(g) ® 2Mg2N(s) D. 4Mg(s) + 3N2(g) ® 2Mg2N3(s) E. 3Mg(s) + N2(g) ® Mg3N2(s)
32. Write a balanced equation for the reaction of calcium hydride with water. A. CaH2(s) + H2O(l) ® CaOH2(s) + H2(g) B. CaH2(s) + H2O(l) ® CaO(s) + 2H2(g) C. CaH2(s) + H2O(l) ® Ca(s) + 2H2(g) + H2O(l) D. CaH2(s) + 2H2O(l) ® Ca(s) + 3H2(g) + O2(l) E. CaH2(s) + 2H2O(l) ® Ca2+(aq) + 2OH-(aq) + 2H2(g)
33. Strontium oxide reacts with water. Write a balanced equation for this reaction. A. SrO(s) + H2O(l) ® Sr2+(aq) + O2-(aq) + H2O(l) B. SrO(s) + H2O(l) ® Sr2+(aq) + 2OH-(aq) C. SrO(s) + H2O(l) ® SrH2(aq) + O2(g) D. SrO(s) + H2O(l) ® Sr2+(aq) + 2H-(aq) + O2(g) E. SrO(s) + H2O(l) ® Sr(s) + H2O2(aq)
34. One product of the reaction of sodium and water is sodium hydroxide. What mass of sodium is needed to make 2.0 L of 0.075 M NaOH(aq)? A. 0.26 g B. 3.4 g C. 6.0 g D. 6.8 g E. 12 g
35. What is the percent composition of chalcopyrite, CuFeS2? A. 25.0% Cu, 25.0% Fe, 50.0% S B. 33.3% Cu, 33.3% Fe, 33.3% S C. 34.6% Cu, 30.4% Fe, 35.0% S D. 41.2% Cu, 17.3% Fe, 41.5% S E. 42.0% Cu, 36.9% Fe, 21.1% S
36. Potassium superoxide is used in self-contained breathing devices. Its function is to A. remove nitrogen from exhaled air. B. absorb smoke or other hazardous vapors. C. generate oxygen only. D. remove exhaled carbon dioxide only. E. both generate oxygen and remove exhaled carbon dioxide.
37. Write a balanced equation for the oxidation of gold by aqua regia. A. Au(s) + 4H+(aq) + 4Cl-(aq) + NO3-(aq) ® AuCl4-(aq) + NO(g) + 2H2O(l) B. Au(s) + 4H+(aq) + NO3-(aq) ® Au3+(aq) + NO(g) + 4H2O(l) C. 2Au(s) + 6H2O(l) ® 2Au(OH)3(s) + 3H2(g) D. 2Au(s) + 6H+(aq) + 6NO3-(aq) ® 2Au3+(aq) + NO3-(aq) + 3H2(g) E. 4Au(s) + 8CN-(aq) + O2(g) + 2H2O(l) ® 4Au(CN)2-(aq) + 4OH-(aq)
38. Write the balanced half-reaction for the reduction of permanganate ion to Mn2+ in an acidic solution. A. MnO4-(aq) + 8H+(aq) ® Mn2+(aq) + 4H2O(l) B. MnO4-(aq) + 5e- ® Mn2+(aq) + 2O2(g) C. MnO4-(aq) + 4H+(aq) + 3e- ® Mn2+(aq) + 4OH-(aq) D. MnO4-(aq) + 4H+(aq) + 5e- ® Mn2+(aq) + 4OH-(aq) E. MnO4-(aq) + 8H+(aq) + 5e- ® Mn2+(aq) + 4H2O(l)
39. Write a balanced equation for the oxidation of copper metal by nitric acid. A. Cu(s) + 6H+(aq) + NO3-(aq) ® Cu2+(aq) + NH3(aq) + 3H2O(l) B. Cu(s) + 12H+(aq) + 2NO3-(aq) ® Cu2+(aq) + N2(g) + 6H2O(l) C. 3Cu(s) + 8H+(aq) + 2NO3-(aq) ® 3Cu2+(aq) + 2NO(g) + 4H2O(l) D. 4Cu(s) + 6H+(aq) + NO3-(aq) ® 4Cu2+(aq) + NH3(aq) + 3H2O(l) E. 5Cu(s) + 12H+(aq) + 2NO3-(aq) ® 5Cu2+(aq) + N2(g) + 6H2O(l)
40. Predict which of the following oxidation-reduction reactions will occur, given the half-reactions and values in the table below. 4H+(aq) + NO3-(aq) + 3e- ® NO(g) + 2H2O(l) = +0.964 V Fe3+(aq) + e- ® Fe2+(s) = +0.771 V 2H+(aq) + 2e- ® H2(g) = 0.000 V Fe2+(aq) + 2e- ® Fe(s) = -0.409 V
1. 2. 3.
3Fe2+(aq) ® Fe(s) + 2Fe3+(aq) Fe(s) + 2H+(aq) ® Fe2+(aq) + H2(g) 3Fe2+(s) + 4H+(aq) + NO3-(aq) ® 3Fe3+(aq) + NO(g) + 2H 2O(l)
A. 1 only B. 2 only C. 3 only D. 2 and 3 E. 1, 2, and 3
41. Given the half-reactions and agent.
values in the table below, determine which species is the best oxidizing
MnO4-(aq) + 8H+(aq) + 5e- ® Mn2+(aq) + 4H2O(l) = +1.512 V 4H+(aq) + NO3-(aq) + 3e- ® NO(g) + 2H2O(l) = +0.964 V Fe3+(aq) + e- ® Fe2+(s) = +0.771 V 2H+(aq) + 2e- ® H2(g) = 0.000 V
A. H+ B. Fe2+ C. Fe3+ D. NO3E. MnO4-
42. Given the half-reactions and
values below, determine which species is the best reducing agent.
MnO4-(aq) + 8H+(aq) + 5e- ® Mn2+(aq) + 4H2O(l) = +1.512 V 4H+(aq) + NO3-(aq) + 3e- ® NO(g) + 2H2O(l) = +0.964 V Fe3+(aq) + e- ® Fe2+(s) = +0.771 V 2H+(aq) + 2e- ® H2(g) = 0.000 V
A. H+ B. H2 C. Fe2+ D. NO3E. MnO443. In an aqueous solution, Cu+ undergoes a reaction to form Cu(s) and Cu2+. This reaction is an example of A. reduction. B. oxidation. C. disproportionation. D. precipitation. E. amphoterism.
44. Even in basic solution, MnO4- can oxidize water. One product is manganese(IV) oxide. Write a balanced chemical equation for the reaction. A. MnO4-(aq) + H2O(l) ® MnO2(s) + H2(g) + OH-(aq) B. MnO4-(aq) + 6H2O(l) ® MnO2(s) + 2H2(g) + 8OH-(aq) C. 2MnO4-(aq) + 2H2O(l) ® 2Mn2+(aq) + 3O2(g) + 4OH-(aq) D. 4MnO4-(aq) + 2H2O(l) ® 4MnO2(s) + 3O2(g) + 4OH-(aq) E. 4MnO4-(aq) + H2O(l) ® 4MnO(s) + O2(g) + 2OH-(aq)
45. Ammonium dichromate, (NH4)2Cr2O7 decomposes in a reaction that resembles the eruption of a volcano. Write a balanced chemical equation for the reaction. A. (NH4)2Cr2O7(s) ® N2(g) + 4H2O(g) + Cr2O3(s) B. (NH4)2Cr2O7(s) ® 2NH3(g) + H2Cr2O7(s) C. (NH4)2Cr2O7(s) ® 2NH3(g) + 4H2O(g) + 2CrO3 D. (NH4)2Cr2O7(s) ® 2NH4+(g) + Cr2O72-(g) E. (NH4)2Cr2O7(s) ® 2NH3(g) + 3H2(g) + H2O(g) + 2CrO3(s)
46. Potassium superoxide is used to generate oxygen in confined places, such as submarines and spacecraft. 4KO2(s) + 2CO2(g) ® K2CO3(s) + 3O2(g) What mass of oxygen will be produced for each 1.0 g KO2 consumed? A. 0.23 g B. 0.34 g C. 0.45 g D. 0.60 g E. 1.0 g
47. Sodium azide is used in air bags for automobiles. The gas produced from the decomposition of sodium azide is nitrogen. 2NaN3(s) ® 2Na(s) + 3N2(g) What mass of NaN3 is required to inflate a bag with a volume of 45 L at a pressure of 1.0 atm and a temperature of 305 K? The gas law constant is 0.0821 Latm/molK and the molar mass of NaN3 is 65.0 g/mol. A. 2.8 ´ 10-2 g B. 1.2 ´ 10-1 g C. 7.8 ´ 101 g D. 1.2 ´ 102 g E. 1.8 ´ 102 g
48. Ilmenite, an ore used in the production of metallic titanium, is composed of 36.8% iron, 31.6% titanium, and 31.6% oxygen. What is the empirical (simplest) formula for ilmenite? A. FeTiO B. FeTiO2 C. FeTiO3 D. Fe2TiO4 E. Fe3Ti2O6
49. Which metal cation is a component of over 70 enzymes? A. Al3+ B. Ca2+ C. Co3+ D. Fe2+ E. Zn2+
50. Which of the following cations is a component of vitamin B12? A. Li+ B. Mg2+ C. Ca2+ D. Co2+ E. Hg2+
51. Which of the following cations is not a trace species in the human body? A. Be2+ B. Cr3+ C. Co2+ D. Zn2+ E. Cu2+
52. Which of the following is the principal cation within cell fluid? A. Na+ B. Mg2+ C. K+ D. Ca2+ E. Fe2+
53. Which of the following cations is not essential to human nutrition? A. Hg2+ B. Zn2+ C. K+ D. Ca2+ E. Fe2+
54. Which of the following cations has the highest concentration in the human body? A. Na+ B. Zn2+ C. K+ D. Ca2+ E. Fe2+
55. What group of elements tend to form colored salts? A. group 1 metals B. group 2 metals C. transition metals D. all of the above E. none of the above
Chapter 20--Chemistry of the Metals Key
1. An ore is A. a natural source from which a metal can be extracted. B. a soluble mixture of ionic compounds. C. a mixture of two or more metallic elements. D. another name for a metal alloy, such as brass. E. a molten mixture of ionic compounds.
2. The principal ores of group 2 metals are A. carbonates. B. chlorides. C. oxides. D. phosphates. E. sulfides.
3. The principal ores of copper, silver and mercury are A. carbonates. B. chlorides. C. oxides. D. phosphates. E. sulfides.
4. Aluminum is produced by electrolysis of a mixture of bauxite ore (Al2O3) and cryolite (Na3AlF6). The purpose of the cryolite is to A. increase the current. B. decrease the potential at which aluminum is reduced. C. convert oxide ion into water D. lower the melting point. E. decrease the pH of the molten solution.
5. The main group metals in Group 1 are referred to as the A. alkali metals. B. alkaline earth metals. C. highly reactive metals. D. transition metals. E. basic metals.
6. The main group metals in Group 2 are referred to as the A. alkali metals. B. alkaline earth metals. C. less reactive metals. D. divalent metals. E. basic metals.
7. Write a balanced equation for the reduction of Co2O3 by carbon monoxide. A. Co2O3(s) + CO(g) ® 2Co(s) + CO4(g) B. 2Co2O3(s) + 2CO(g) ® 2Co2CO3(s) +O2(g) C. Co2O3(s) + CO(g) ® Co(s) + CO2(g) D. Co2O3(s) + CO(g) ® Co(s) + C(g) + 2O2(g) E. Co2O3(s) + 3CO(g) ® 2Co(s) + 3CO2(g)
8. Write a balanced equation for the reduction of manganese(IV) oxide ore by carbon monoxide. A. MnO(s) + CO(g) ® Mn(s) + CO2(g) B. MnO(s) + 2CO(g) ® Mn(s) + 2CO2(g) C. Mn2O(s) + CO(g) ® Mn(s) + CO2(g) D. MnO2(s) + 2CO(g) ® Mn(s) + 2CO2(g) E. MnO2(s) + 2CO(g) ® Mn(s) + 2C(s) + O2(g)
9. Both sodium metal and chlorine gas are created by the electrolysis of molten sodium chloride. If 1.00 ´ 102 amperes of current is passed through the cathode for 8.00 hr, what mass of Cl2 will be formed? (F = 96480 C/mol e-). A. 0.588 g B. 1.18 g C. 17.6 g D. 1.06 ´ 103 g E. 2.12 ´ 102 g
10. What is a useful by-product of the production of sodium metal? A. hydrogen B. magnesium C. chlorine D. helium E. calcium
11. Who worked out the process for obtaining aluminum from bauxite ore? A. Michael Faraday B. Enrico Fermi C. Charles Hall D. Linus Pauling E. Alfred Werner
12. Aluminum is obtained by A. the electrolysis of cryolite, Na3AlF6. B. reduction by a strong inorganic reducing agent, such as thiosulfate ion, S2O32-. C. the reduction of Al2O3 by carbon monoxide in a blast furnace. D. chemical decomposition of Al2O3 at temperatures greater than 2000C. E. electrolysis of a molten mixture of bauxite ore and cryolite, Na3AlF6.
13. What is the most common chemical reducing agent in metallurgical processes? A. thiosulfate ion, S2O32B. oxygen, O2 C. sodium metal D. carbon; in the form of carbon monoxide E. carbon; in the form of carbon dioxide
14. Write a balanced chemical equation for the chemical reduction of iron(III) oxide that occurs in a blast furnace. A. Fe2O3(s) + 3CO(g) ® 2Fe(s) + 3CO2(g) B. Fe2O3(s) + 3CO2(g) ® 2Fe(s) + 3CO(g) C. Fe2O3(s) + 3C(g) ® 2Fe(s) + 3CO(g) D. FeO(s) + CO2(g) ® FeCO3(s) E. FeO(s) + CO(g) ® Fe(s) + CO2(g)
15. At temperatures greater than 800C, limestone decomposes according to which balanced chemical equation? A. CaO(s) ® Ca(s) + O(g) B. 2CaO(s) ® 2Ca(s) + O2(g) C. CaCO3(s) ® CaO(s) + CO2(g) D. CaCO3(s) ® CaO2(s) + CO(g) E. CaCO2(g) ® Ca(s) + CO2(g)
16. Slag is a glassy material that forms when calcium oxide reacts with impurities (mostly silicon dioxide) in a blast furnace. Write the balanced chemical equation for the reaction. A. CaCO3(s) + Si(s) ® CaSiO3(l) B. CaCO3(s) + SiO2(s) ® CaSiO5(l) C. CaO(s) + SiO2(s) ® CaSiO(l) + O2(g) D. CaO(s) + SiO2(s) ® CaSiO3(l) E. CaO(s) + SiO(s) ® CaCO2(l)
17. Iron that is made in a blast furnace is called pig iron. The pig iron is converted into steel by reducing its carbon content from roughly 4% to under 2% by A. reduction of the carbon with hydrogen gas. B. the addition of calcium oxide to form slag. C. vacuum filtration. D. electrolysis. E. reaction with oxygen to form CO2(g).
18. The iron produced in a blast furnace contains significant impurities of carbon, silicon, manganese, and phosphorus. These impurities are removed in a ________. A. basic oxygen furnace B. distillation column C. mass spectrometer D. electrolytic cell E. centrifuge
19. Copper can be obtained from copper sulfide by heating in the presence of oxygen. This process is called A. roasting. B. oxidation. C. smelting. D. steaming. E. alloying.
20. The roasting of Cu2S(s) produces Cu(s) and SO2(g). What mass of sulfur dioxide is produced by the roasting of 1.00 ´ 103 kg Cu2S(s)? Cu2S(s) + O2(g) ® 2Cu(s) + SO2(g) A. 201 B. 402 kg C. 798 kg D. 804 kg E. 2.48 ´ 103 kg
21. Write a balanced chemical equation for the reaction of HgS with O2. A. HgS(s) + O2(g) ® HgO2(s) + S(s) B. HgS(s) + O2(g) ® Hg(l) + SO2(g) C. HgS(s) + O2(g) ® HgSO2(s) D. HgS(s) + 2O2(g) ® HgSO4(s) E. 2HgS(s) + O2(g) ® HgO(s) + 2S(s)
22. All of the following metals may be found in nature as free elements EXCEPT ____. A. Ir B. Au C. Pt D. Ti E. Rh
23. Separation of gold from ore is currently accomplished by using which ligand to make a soluble coordination compound from Au+1? A. chloride ion B. cyanide ion C. ammonia D. perchlorate ion E. ethylenediamine
24. Historically, miners separated gold by taking advantage of its A. high density. B. low melting point. C. high solubility in water. D. radioactivity. E. amphoteric nature.
25. Write a balanced equation for the reaction of sodium and bromine. A. Na(s) + Br-(l) ® NaBr(s) B. Na(s) + Br(l) ® NaBr(s) + Br-(aq) C. Na(s) + Br2(l) ® NaBr2(s) D. 2Na(s) + Br2(l) ® 2NaBr(s) E. 4Na(s)) + Br2(l) ® 2Na2Br(s)
26. Write a balanced chemical equation for the reaction of potassium and water. A. K(s) + H2O(l) ® KO(s) + H2(g) B. 2K(s) + 2H2O(l) ® 2KOH(aq) + H2(g) C. 2K(s) + 2H2O(l) ® K2O2(s) + 2H2(g) D. 2K(s) + H2O(l) ® K2O(s) + H2(g) E. 4K(s) + 2H2O(l) ® 4KH(s) + O2(g)
27. Write a balanced equation for the reaction of barium and nitrogen gas. A. Ba(s) + N2(g) ® BaN2(s) B. 2Ba(s) + N2(g) ® 2BaN(s) C. 2Ba(s) + N2(g) ® 2BaN(s) D. 3Ba(s) + N2(g) ® Ba3N2(s) E. 4Ba(s) + 3N2(g) ® 2Ba2N3(s)
28. Which of the following metals will react most violently with water? A. Li B. Mg C. K D. Ca E. Cr
29. Write a balanced equation for the reaction of Na and O2 to form sodium peroxide. A. 4Na(s) + O2(g) ® 2Na2O(s) B. Na(s) + O2(g) ® NaO2(s) C. 2Na(s) + O2(g) ® Na2O2(s) D. 2Na(s) + O2(g) ® 2NaO(s) E. 4Na(s) + 3O2(g) ® 2Na2O3(s)
30. The principal product of the reaction of potassium and oxygen is not K2O. The product is named ____ and has the formula ____. A. dipotassium trioxide, K2O3 B. potassium monoxide, KO C. potassium superoxide, KO2 D. potassium peroxide, K2O2 E. potassium trioxide, KO3
31. Write a balanced equation for the reaction of magnesium and nitrogen that occurs upon heating. A. Mg(s) + N2(g) ® MgN2(s) B. 2Mg(s) + N2(g) ® 2MgN(s) C. 4Mg(s) + N2(g) ® 2Mg2N(s) D. 4Mg(s) + 3N2(g) ® 2Mg2N3(s) E. 3Mg(s) + N2(g) ® Mg3N2(s)
32. Write a balanced equation for the reaction of calcium hydride with water. A. CaH2(s) + H2O(l) ® CaOH2(s) + H2(g) B. CaH2(s) + H2O(l) ® CaO(s) + 2H2(g) C. CaH2(s) + H2O(l) ® Ca(s) + 2H2(g) + H2O(l) D. CaH2(s) + 2H2O(l) ® Ca(s) + 3H2(g) + O2(l) E. CaH2(s) + 2H2O(l) ® Ca2+(aq) + 2OH-(aq) + 2H2(g)
33. Strontium oxide reacts with water. Write a balanced equation for this reaction. A. SrO(s) + H2O(l) ® Sr2+(aq) + O2-(aq) + H2O(l) B. SrO(s) + H2O(l) ® Sr2+(aq) + 2OH-(aq) C. SrO(s) + H2O(l) ® SrH2(aq) + O2(g) D. SrO(s) + H2O(l) ® Sr2+(aq) + 2H-(aq) + O2(g) E. SrO(s) + H2O(l) ® Sr(s) + H2O2(aq)
34. One product of the reaction of sodium and water is sodium hydroxide. What mass of sodium is needed to make 2.0 L of 0.075 M NaOH(aq)? A. 0.26 g B. 3.4 g C. 6.0 g D. 6.8 g E. 12 g
35. What is the percent composition of chalcopyrite, CuFeS2? A. 25.0% Cu, 25.0% Fe, 50.0% S B. 33.3% Cu, 33.3% Fe, 33.3% S C. 34.6% Cu, 30.4% Fe, 35.0% S D. 41.2% Cu, 17.3% Fe, 41.5% S E. 42.0% Cu, 36.9% Fe, 21.1% S
36. Potassium superoxide is used in self-contained breathing devices. Its function is to A. remove nitrogen from exhaled air. B. absorb smoke or other hazardous vapors. C. generate oxygen only. D. remove exhaled carbon dioxide only. E. both generate oxygen and remove exhaled carbon dioxide.
37. Write a balanced equation for the oxidation of gold by aqua regia. A. Au(s) + 4H+(aq) + 4Cl-(aq) + NO3-(aq) ® AuCl4-(aq) + NO(g) + 2H2O(l) B. Au(s) + 4H+(aq) + NO3-(aq) ® Au3+(aq) + NO(g) + 4H2O(l) C. 2Au(s) + 6H2O(l) ® 2Au(OH)3(s) + 3H2(g) D. 2Au(s) + 6H+(aq) + 6NO3-(aq) ® 2Au3+(aq) + NO3-(aq) + 3H2(g) E. 4Au(s) + 8CN-(aq) + O2(g) + 2H2O(l) ® 4Au(CN)2-(aq) + 4OH-(aq)
38. Write the balanced half-reaction for the reduction of permanganate ion to Mn2+ in an acidic solution. A. MnO4-(aq) + 8H+(aq) ® Mn2+(aq) + 4H2O(l) B. MnO4-(aq) + 5e- ® Mn2+(aq) + 2O2(g) C. MnO4-(aq) + 4H+(aq) + 3e- ® Mn2+(aq) + 4OH-(aq) D. MnO4-(aq) + 4H+(aq) + 5e- ® Mn2+(aq) + 4OH-(aq) E. MnO4-(aq) + 8H+(aq) + 5e- ® Mn2+(aq) + 4H2O(l)
39. Write a balanced equation for the oxidation of copper metal by nitric acid. A. Cu(s) + 6H+(aq) + NO3-(aq) ® Cu2+(aq) + NH3(aq) + 3H2O(l) B. Cu(s) + 12H+(aq) + 2NO3-(aq) ® Cu2+(aq) + N2(g) + 6H2O(l) C. 3Cu(s) + 8H+(aq) + 2NO3-(aq) ® 3Cu2+(aq) + 2NO(g) + 4H2O(l) D. 4Cu(s) + 6H+(aq) + NO3-(aq) ® 4Cu2+(aq) + NH3(aq) + 3H2O(l) E. 5Cu(s) + 12H+(aq) + 2NO3-(aq) ® 5Cu2+(aq) + N2(g) + 6H2O(l)
40. Predict which of the following oxidation-reduction reactions will occur, given the half-reactions and values in the table below. 4H+(aq) + NO3-(aq) + 3e- ® NO(g) + 2H2O(l) = +0.964 V Fe3+(aq) + e- ® Fe2+(s) = +0.771 V 2H+(aq) + 2e- ® H2(g) = 0.000 V Fe2+(aq) + 2e- ® Fe(s) = -0.409 V
1. 2. 3.
3Fe2+(aq) ® Fe(s) + 2Fe3+(aq) Fe(s) + 2H+(aq) ® Fe2+(aq) + H2(g) 3Fe2+(s) + 4H+(aq) + NO3-(aq) ® 3Fe3+(aq) + NO(g) + 2H 2O(l)
A. 1 only B. 2 only C. 3 only D. 2 and 3 E. 1, 2, and 3
41. Given the half-reactions and agent.
values in the table below, determine which species is the best oxidizing
MnO4-(aq) + 8H+(aq) + 5e- ® Mn2+(aq) + 4H2O(l) = +1.512 V 4H+(aq) + NO3-(aq) + 3e- ® NO(g) + 2H2O(l) = +0.964 V Fe3+(aq) + e- ® Fe2+(s) = +0.771 V 2H+(aq) + 2e- ® H2(g) = 0.000 V
A. H+ B. Fe2+ C. Fe3+ D. NO3E. MnO4-
42. Given the half-reactions and
values below, determine which species is the best reducing agent.
MnO4-(aq) + 8H+(aq) + 5e- ® Mn2+(aq) + 4H2O(l) = +1.512 V 4H+(aq) + NO3-(aq) + 3e- ® NO(g) + 2H2O(l) = +0.964 V Fe3+(aq) + e- ® Fe2+(s) = +0.771 V 2H+(aq) + 2e- ® H2(g) = 0.000 V
A. H+ B. H2 C. Fe2+ D. NO3E. MnO443. In an aqueous solution, Cu+ undergoes a reaction to form Cu(s) and Cu2+. This reaction is an example of A. reduction. B. oxidation. C. disproportionation. D. precipitation. E. amphoterism.
44. Even in basic solution, MnO4- can oxidize water. One product is manganese(IV) oxide. Write a balanced chemical equation for the reaction. A. MnO4-(aq) + H2O(l) ® MnO2(s) + H2(g) + OH-(aq) B. MnO4-(aq) + 6H2O(l) ® MnO2(s) + 2H2(g) + 8OH-(aq) C. 2MnO4-(aq) + 2H2O(l) ® 2Mn2+(aq) + 3O2(g) + 4OH-(aq) D. 4MnO4-(aq) + 2H2O(l) ® 4MnO2(s) + 3O2(g) + 4OH-(aq) E. 4MnO4-(aq) + H2O(l) ® 4MnO(s) + O2(g) + 2OH-(aq)
45. Ammonium dichromate, (NH4)2Cr2O7 decomposes in a reaction that resembles the eruption of a volcano. Write a balanced chemical equation for the reaction. A. (NH4)2Cr2O7(s) ® N2(g) + 4H2O(g) + Cr2O3(s) B. (NH4)2Cr2O7(s) ® 2NH3(g) + H2Cr2O7(s) C. (NH4)2Cr2O7(s) ® 2NH3(g) + 4H2O(g) + 2CrO3 D. (NH4)2Cr2O7(s) ® 2NH4+(g) + Cr2O72-(g) E. (NH4)2Cr2O7(s) ® 2NH3(g) + 3H2(g) + H2O(g) + 2CrO3(s)
46. Potassium superoxide is used to generate oxygen in confined places, such as submarines and spacecraft. 4KO2(s) + 2CO2(g) ® K2CO3(s) + 3O2(g) What mass of oxygen will be produced for each 1.0 g KO2 consumed? A. 0.23 g B. 0.34 g C. 0.45 g D. 0.60 g E. 1.0 g
47. Sodium azide is used in air bags for automobiles. The gas produced from the decomposition of sodium azide is nitrogen. 2NaN3(s) ® 2Na(s) + 3N2(g) What mass of NaN3 is required to inflate a bag with a volume of 45 L at a pressure of 1.0 atm and a temperature of 305 K? The gas law constant is 0.0821 Latm/molK and the molar mass of NaN3 is 65.0 g/mol. A. 2.8 ´ 10-2 g B. 1.2 ´ 10-1 g C. 7.8 ´ 101 g D. 1.2 ´ 102 g E. 1.8 ´ 102 g
48. Ilmenite, an ore used in the production of metallic titanium, is composed of 36.8% iron, 31.6% titanium, and 31.6% oxygen. What is the empirical (simplest) formula for ilmenite? A. FeTiO B. FeTiO2 C. FeTiO3 D. Fe2TiO4 E. Fe3Ti2O6
49. Which metal cation is a component of over 70 enzymes? A. Al3+ B. Ca2+ C. Co3+ D. Fe2+ E. Zn2+
50. Which of the following cations is a component of vitamin B12? A. Li+ B. Mg2+ C. Ca2+ D. Co2+ E. Hg2+
51. Which of the following cations is not a trace species in the human body? A. Be2+ B. Cr3+ C. Co2+ D. Zn2+ E. Cu2+
52. Which of the following is the principal cation within cell fluid? A. Na+ B. Mg2+ C. K+ D. Ca2+ E. Fe2+
53. Which of the following cations is not essential to human nutrition? A. Hg2+ B. Zn2+ C. K+ D. Ca2+ E. Fe2+
54. Which of the following cations has the highest concentration in the human body? A. Na+ B. Zn2+ C. K+ D. Ca2+ E. Fe2+
55. What group of elements tend to form colored salts? A. group 1 metals B. group 2 metals C. transition metals D. all of the above E. none of the above
Chapter 21--Chemistry of the Nonmetals 1. Which of the following elements is the strongest oxidizing agent? A. N2 B. O2 C. F2 D. Cl2 E. P4
2. Which of the following elements is the least chemically reactive? A. N2 B. O2 C. F2 D. Cl2 E. P4
3. Which of the following molecules is a product of most chemical explosives? A. H2 B. N2 C. O2 D. SO2 E. HCl
4. Which of the halogens are solids at 298 K and 1 atmosphere? A. Cl2 B. Br2 C. I2 D. Cl2 and Br2 E. Br2 and I2
5. What intermolecular bonds or forces are present in P4(s)? A. hydrogen bonds B. covalent bonds C. ionic bonds D. dipole forces E. dispersion forces
6. Which of the following reactions represents the disproportionation of chlorine in water? A. Cl2(g) + H2O(l) Cl-(aq) + H2(g) + ClO+(aq) B. Cl2(g) + H2O(l) Cl-(aq) + H+(aq) + HClO(aq) C. Cl2(g) + H2O(l) Cl-(aq) + HCl(aq) + OH-(aq) D. Cl2(g) + 2H2O(l) Cl-(aq) + 3H+(aq) + HClO2(aq) E. Cl2(g) + 4H2O(l) Cl-(aq) + 8H+(aq) + ClO4-(aq)
7. What is the molecular formula for chlorous acid? A. HCl B. HClO C. H2ClO D. HClO2 E. HClO3
8. Which of the following lists of nonmetals contains only elements that can be found in nature in their elemental forms? A. chlorine, nitrogen, oxygen B. silicon, arsenic, sulfur C. nitrogen, phosphorus, arsenic D. bromine, carbon, oxygen E. nitrogen, oxygen, sulfur
9. Fluorine and chlorine gases are primarily produced by A. distillation. B. oxidation by electrolysis. C. oxidation by hydronium ion. D. reduction by carbon monoxide. E. roasting with copper(II) sulfide.
10. Write a balanced net ionic equation for the reaction of chlorine gas with aqueous potassium iodide. A. Cl2(g) + 2KI(aq) ® 2KCl(s) + I2(aq) B. Cl2(g) + 2I-(aq) ® 2ICl(aq) C. Cl2(g) + 2I-(aq) ® 2Cl-(aq) + I2(aq) D. Cl2(g) + 2I-(aq) ® Cl2I(aq) + I-(aq) E. Cl2(g) + K+(aq) + I-(aq) ® KCl(s) + ICl(g)
11. Which of the following statements are true about the reaction below? NH3 + BF3 ® NH3BF3 1. 2. 3.
Ammonia acts as a Lewis base. Ammonia acts as a Brønsted-Lowry base. Ammonia acts as a reducing agent.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 2 and 3 12. Which of the following nitrogen containing compounds or ions is not capable of acting as an oxidizing agent? A. NO3B. NO2C. NH3 D. N2 E. NO2
13. Hydrazine is made commercially by reacting aqueous ammonia with hypochlorite ion. Write a balanced equation for this reaction. A. 2NH3(aq) + ClO4-(aq) ® N2H4(aq) + ClO3-(aq) + H2O(l) B. 2NH3(aq) + ClO3-(aq) ® N2H4(aq) + ClO2-(aq) + H2O(l) C. 2NH3(aq) + ClO2-(aq) ® N2H4(aq) + Cl-(aq) + H2O2(aq) D. 2NH3(aq) + ClO-(aq) ® N2H4(aq) + Cl-(aq) + H2O(l) E. 2NH3(aq) + ClO-(aq) ® N2H4(aq) + H2(aq) + ClO-(aq)
14. Write a balanced equation for the reaction of gaseous ammonia with hydrogen chloride. A. NH3(g) + HCl(g) ® NH2Cl(g) + H2(g) B. NH3(g) + HCl(g) ® NH4Cl(s) C. NH3(g) + 3HCl(g) ® NCl3(g) + 3H2(g) D. NH3(g) + 2HCl(g) ® NH5(g) + Cl2(g) E. 2NH3(g) + 2HCl(g) ® N2(g) + 4H2(g) + Cl2(g)
15. Which of the following statements best describe the role of hydrogen sulfide in the reaction below? 2H2S(aq) + O2(g) ® 2S(s) + 2H2O(l) 1. 2. 3.
H2S is a reducing agent. H2S is a Brønsted-Lowry base. H2S is a precipitating agent.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 2 and 3 16. Write a half-reaction for the reduction of aqueous hydrogen peroxide in an acidic solution. A. H2O2(aq) + 2e- ® H2(g) + O22-(aq) B. 2H2O2(aq) + 2e- ® O2(g) + 2H2O(l) C. H2O2(aq) + 2e- ® 2OH-(aq) D. H2O2(aq) + 2H+(aq) + 2e- ® 2H2O(l) E. Hydrogen peroxide cannot be reduced under acidic conditions.
17. Write a half-reaction for the oxidation of hydrogen peroxide. A. H2O2(aq) ® 2OH-(aq) + 2eB. H2O2(aq) ® H2(g) + O2(g) C. H2O2(aq) ® H2(g) + O2(g) + 2eD. 2H2O2(aq) ® O2(g) + 2H2O(l) + 2eE. H2O2(aq) ® O2(g) + 2H+(aq) + 2e-
18. What is the overall reaction for the disproportionation of hydrogen peroxide in water? A. H2O2(aq) ® 2OH-(aq) B. H2O2(aq) ® O2(g) + H2(g) C. 2H2O2(aq) ® O2(g) + 2H2O(l) D. H2O2(aq) ® O2(g) + 2H+(aq) E. H2O2(aq) + H2O(l) ® 4H+(aq) + O3(g)
19. Write a balanced chemical equation for the reaction of hydrogen peroxide and iodide ion in an acidic solution. A. H2O2(aq) + 2I-(aq) + 2H+(aq) ® I2(s) + 2H2O(l) B. H2O2(aq) + 2I-(aq) ® I2(s) + 2OH-(aq) C. H2O2(aq) + I2(s) ® 2I-(aq) + H2(g) + O2(g) D. H2O2(aq) + I2(s) + 2H+(aq) ® 2HI(s) + 2OH-(aq) E. H2O2(aq) + 2I-(aq) ® I2(s) + H2(g) + O2(g)
20. Write a balanced chemical equation for the reaction of aqueous solutions of hydrofluoric acid and potassium carbonate. A. 2HF(aq) + CO32-(aq) ® CO2(g) + H2O(l) + 2F-(aq) B. 2HF(aq) + K2CO3(aq) ® F2CO3(s) + H2(g) C. 2HF(aq) + K2CO3(s) ® 2KF(s) + H2CO3(s) D. 2H+(aq) + CO32-(aq) ® H2CO3(s) E. 2H+(aq) + CO32-(aq) ® CO2(g) + H2O(l)
21. Which of the following acids is never stored in glass containers because it reacts with silicon dioxide and ionic silicates? A. HNO3 B. HClO4 C. H2SO4 D. HF E. HI
22. All of the following oxidation states are observed for sulfur EXCEPT _____. A. -2 B. +2 C. +4 D. +6 E. +7
23. Which reddish-brown gas is a major factor in the formation of photochemical smog? A. NO B. NO2 C. SO2 D. CO2 E. CO
24. Which of the following molecules are paramagnetic: N2O, NO, and NO2? A. N2O only B. NO only C. NO2 only D. N2O and NO E. NO and NO2
25. Which of the following molecules has the ability to dilate blood vessels, thereby lowering blood pressure? A. NO B. N2O C. CO D. HCl E. SO3
26. Write a balanced chemical equation for the reaction of N2O5 and water. A. N2O5(s) + H2O(l) ® N2O4(s) + H2O2(l) B. N2O5(s) + H2O(l) ® N2O(g) + H2O(l) + 2O2(g) C. N2O5(s) + 3H2O(l) ® NH3(g) + 4O2(g) D. N2O5(s) + H2O(l) ® 2HNO3(l) E. N2O5(s) + H2O(l) ® 2HNO2(l) + O2(g)
27. Write a balanced chemical equation for the combustion of white phosphorus. A. P4(s) + 8O2(g) ® 4PO4(g) B. P4(s) + 3O2(g) ® 2P2O3(s) C. P4(s) + 5O2(g) ® P4O10(s) D. P4(s) + 8O2(g) ® 4PO43-(g) E. P4(s) + O2(g) ® 2P2O(g)
28. In which of the following compounds does phosphorus have an oxidation number of +3? A. P4S3 B. H3PO3 C. H3PO4 D. P4O10 E. PH3
29. Nonmetal oxides react with water to form A. solids. B. fullerenes. C. Lewis bases. D. Brønsted-Lowry acids. E. precipitates.
30. Write a balanced chemical equation for the reaction of dinitrogen trioxide with water. A. N2O3(g) + H2O(l) ® N2O4(g) + H2(g) B. N2O3(g) + H2O(l) ® 2HNO2(aq) C. N2O3(g) + 3H2O(l) ® 2HNO3(aq) + 2H2(g) D. N2O3(g) + H2O(l) ® N2H2(g) + 2O2(g) E. N2O3(g) + 3H2O(l) ® 2NH3(g) + 3O2(g)
31. Write a balanced chemical equation for the reaction of sulfur trioxide with water. A. SO3(g) + H2O(l) ® S(s) + H2(g) + 2O2(g) B. SO3(g) + H2O(l) ® H2SO3(aq) C. SO3(g) + H2O(l) ® SO2(g) + H2O2(g) D. SO3(g) + H2O(l) ® H2S(aq) + 2O2(g) E. SO3(g) + H2O(l) ® H2SO4(aq)
32. Which of the following oxoacids has the largest Ka value? A. HClO4 B. HClO3 C. HClO2 D. HClO E. All of the above oxoacids have similar Ka values.
33. Place the following acids in order from smallest to largest Ka value: HClO, HBrO, and HIO. A. HClO < HBrO < HIO B. HClO < HIO < HBrO C. HBrO < HClO < HIO D. HBrO < HIO < HClO E. HIO < HBrO < HClO
34. Calculate the pH of a 0.35 M solution of sodium hypobromite, NaClO. The acid dissociation constant for HBrO is 2.6 ´ 10-9. A. 2.94 B. 4.96 C. 8.13 D. 9.04 E. 11.06
35. H3PO4 is a weak, triprotic acid. The acid dissociation constants are Ka1 = 7.1 ´ 10-3, Ka2 = 6.2 ´ 10-8, and Ka3 = 4.5 ´ 10-13. Calculate the equilibrium constant for the reaction below. H3PO4(aq) PO43-(aq) + 3H+(aq) A. 2.0 ´ 10-22 B. 4.4 ´ 10-10 C. 2.0 ´ 10-8 D. 8.7 ´ 10-6 E. 7.1 ´ 10-3
36. Which of the following oxoanions of chlorine cannot act as an oxidizing agent? A. HClO B. HClO2 C. HClO3 D. HClO4 E. All of the above can act as oxidizing agents.
37. Which of the following nitrogen containing compounds is not able to act as an oxidizing agent? A. HNO3 B. NH3 C. HNO2 D. N2H4 E. All of the above can act as oxidizing agents.
38. Write a balanced chemical equation for the disproportionation of chlorate ion in water. The products are chloride and perchlorate ions. A. 2ClO3-(aq) + H2O(l) ® Cl-(aq) + ClO4-(aq) + 2H+(aq) B. 3ClO3-(aq) + 2H2O(l) ® Cl-(aq) + 2ClO4-(aq) + 4H+(aq) C. 4ClO3-(aq) ® Cl-(aq) + 3ClO4-(aq) D. 4ClO3-(aq) ® 2ClO2-(aq) + 2ClO4-(aq) E. 5ClO3-(aq) + 6H+(aq) ® 2Cl-(aq) + 3ClO4-(aq) + 3H2O(l)
39. The standard reduction potential of H2O2 = +1.763 V. Which of the following species will be oxidized by hydrogen peroxide?
Co2+(aq) ® Co3+(aq) + 2e Ag(s) ® Ag+(aq) + eCo(s) ® Co2+(aq) + 2eZn(s) ® Zn2+(aq) + 2eNa(s) ® Na+(aq) + e-
+1.953 V +0.799 V -0.282 V -0.762 V -2.714 V
A. Co2+ B. Co2+ and Ag(s) C. Co(s) and Co2+ D. Zn(s) and Na(s) E. Ag(s), Co(s), Zn(s), and Na(s) 40. The first step in the synthesis of nitric acid involves the reaction of ammonia with oxygen to produce nitrogen monoxide. 4NH3(g) + 5O2(g) ® 4NO(g) + 6H2O(g) What volume of NO is produced if 14.0 L NH3 is reacted with 15.0 L O2? Assume the reaction is 100% efficient. A. 12.0 L B. 14.0 L C. 15.0 L D. 23.2 L E. 29.0 L
41. Write a balanced chemical equation for the oxidation of iron metal with nitric acid. Assume the reaction products are Fe3+ and NO gas. A. Fe(s) + NO3-(aq) ® Fe3+(aq) + NO(g) + O2(g) B. Fe(s) + HNO3(aq) ® Fe3+(aq) + NO(g) + OH-(aq) C. Fe(s) + 4H+(aq) + NO3-(aq) ® Fe3+(aq) + NO(g) + 2H2O(l) D. Fe(s) + H+(aq) + NO3-(aq) ® Fe3+(aq) + NO(g) + O2H-(aq) E. 2Fe(s) + 12H+(aq) + 3NO3-(aq) ® 2Fe3+(aq) + 3NO(g) + 6H2O(l)
42. Write a balanced chemical equation for the oxidation of chromium metal with nitric acid. Assume the reaction products are Cr3+ and NH4+. A. H+(aq) + NO3-(aq) + Cr(s) ® NH4+(aq) + Cr3+(aq) B. 4H+(aq) + NO3-(aq) + Cr(s) ® NH4+(aq) + Cr3+(aq) C. 10H+(aq) + NO3-(aq) + 3Cr(s) ® NH4+(aq) + 3Cr3+(aq) + 3H2O(l) D. 10H+(aq) + NO3-(aq) + Cr(s) ® NH4+(aq) + Cr3+(aq) + 3H2O(l) E. 30H+(aq) + 3NO3-(aq) + 8Cr(s) ® 3NH4+(aq) + 8Cr3+(aq) + 9H2O(l)
43. Which of the following compounds will react with water to produce a basic solution: NH3, H2O2, and H2S? A. NH3 only B. H2O2 only C. H2S only D. NH3 and H2S E. H2O2 and H2S
44. All of the following statements concerning oxoanions and oxoacids are true EXCEPT A. MnO4- can never act as an oxidizing agent. B. Cr2O72- can never act as a reducing agent. C. ClO4- is a stronger oxidizing agent than ClO-. D. ClO2- can behave as either an oxidizing or a reducing agent. E. HClO3 is a stronger acid than HClO2.
45. Nitric acid solutions slowly turn brown when exposed to sunlight. Which of the following reactions is responsible? A. NO3-(aq) ® NO(g) + O2(g) B. 2NO3-(aq) ® 2NO2(g) + O2(g) C. 2H+(aq) + NO3-(aq) ® NO2(g) + H2O(l) D. 4H+(aq) + 4NO3-(aq) ® 4NO2(g) + 2H2O(l) + O2(g) E. 4H+(aq) + NO3-(aq) ® NO(g) + 2H2(g) + O2(g)
46. Hydrazine, N2H4, is produced by the Raschig process-the oxidation of ammonia with alkaline ____. A. HNO3(aq) B. NaOCl(aq) C. N2(g) D. HPO42–(aq) E. N2O(g)
47. Sulfuric acid can act as 1. 2. 3.
a drying agent. an oxidizing agent. a Brønsted-Lowry acid.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 48. Write a balanced equation for the reaction of table sugar, C12H22O11, with concentrated sulfuric acid. A. C12H22O11(s) ® 12C(s) + 11H2O(l) B. C12H22O11(s) + H2SO4(aq) ® C12H23O11+(s) + HSO4-(aq) C. C12H22O11(s) + H2SO4(aq) + 6O2(g) ® 12CO2(g) + SO2(g) + H2O(l) D. C12H22O11(s) + 2H+(aq) + SO42-(aq) ® C12H24O112+(s) + SO42-(aq) E. C12H22O11(s) + 2H+(aq) + SO42-(aq) ® 12CO(g) + SO2(g) + 2OH+(aq) + H2O(l)
49. Which of the following statements is INCORRECT? A. Sodium dihydrogen phosphate, NaH2PO4, is used in acidic cleaners. B. Sodium phosphate, Na3PO4, is used in strongly basic cleaners. C. Calcium dihydrogen phosphate, Ca(H2PO4)2 is used in fertilizers. D. The principal use of Na2HPO4 is as an emulsifier in the manufacture of cheese. E. Phosphoric acid is a strong acid, a powerful oxidizing agent, and a drying agent.
50. Write a balanced equation for the reaction of concentrated sulfuric acid with copper metal. Assume two products are sulfur dioxide and Cu2+. A. Cu(s) + SO42-(aq) ® Cu2+(aq) + O2(g) + SO2(g) B. Cu(s) + 4H+(aq) + SO42-(aq) ® CuS(s) + 2O2(g) C. Cu(s) + 2H+(aq) + SO42-(aq) ® Cu2+(aq) + H2O(l) + SO2(g) D. Cu(s) + 4H+(aq) + SO42-(aq) ® Cu2+(aq) + 2H2O(l) + SO2(g) E. Cu(s) + 8H+(aq) + SO42-(aq) ® Cu2+(aq) + 4H2O(l) + S(g)
51. Write a balanced equation for the reaction of Al3+ with aqueous ammonia. A. Al3+(aq) + 2NH3(aq) ® Al(s) + 3H2(g) + N2(g) B. Al3+(aq) + 3NH3(aq) + 3H2O(l) ® Al(OH)3(s) + 3NH4+(aq) C. Al3+(aq) + 2NH3(aq) ® Al(NH3)23+(aq) D. Al3+(aq) + 3NH3(aq) + 3H+(aq) ® Al(NH4)36+(aq) E. Al3+(aq) + 6NH3(aq) ® Al(NH3)63+(aq)
52. Write a balanced equation for the reaction of Ni2+ with H2S. A. Ni2+(aq) + 2H2S(aq) ® Ni(HS)2(s) + 2H+(aq) B. Ni2+(aq) + H2S(aq) ® NiH2(s) + S2-(aq) C. Ni2+(aq) + H2S(aq) ® NiS(s) + 2H+(aq) D. Ni2+(aq) + H2S(aq) ® Ni(s) + S(s) + 2H+(aq) E. No reaction will occur.
53. Hydrogen sulfide is used to remove sulfur dioxide in power plant smokestacks. The reaction produces elemental sulfur and water. Write a balanced chemical equation for this reaction. A. H2S(l) + SO2(g) ® 2HSO(s) B. H2S(l) + SO2(g) ® 2S(s) + 2OH-(aq) C. 2H2S(l) + SO2(g) ® 3S(s) + 2H2O(l) D. H2S(l) + SO2(g) ® 2S(s) + H2O2(aq) E. H2S(l) + SO2(g) + 2H+(aq) ® 2S(s) + 2H2O(l)
54. In modern forensic laboratories, arsenic poisoning is detected by testing A. a strand of hair. B. samples of brain tissue. C. samples of liver tissue. D. heart tissue. E. bone fragments.
55. As one might predict, one allotrope of arsenic has the formula A. As4 B. As8 C. As2 D. As3 E. As10
56. The average concentration of bromine (as bromide) in seawater is 65 ppm. Calculate the volume of seawater (d = 64.0 lb/ft3) in cubic feet required to produce one kilogram of liquid bromine. A. 109 ft3 B. 448 ´ 103 ft3 C. 240 ft3 D. 154 ´ 102 ft3 E. 529 ft3
Chapter 21--Chemistry of the Nonmetals Key
1. Which of the following elements is the strongest oxidizing agent? A. N2 B. O2 C. F2 D. Cl2 E. P4
2. Which of the following elements is the least chemically reactive? A. N2 B. O2 C. F2 D. Cl2 E. P4
3. Which of the following molecules is a product of most chemical explosives? A. H2 B. N2 C. O2 D. SO2 E. HCl
4. Which of the halogens are solids at 298 K and 1 atmosphere? A. Cl2 B. Br2 C. I2 D. Cl2 and Br2 E. Br2 and I2
5. What intermolecular bonds or forces are present in P4(s)? A. hydrogen bonds B. covalent bonds C. ionic bonds D. dipole forces E. dispersion forces
6. Which of the following reactions represents the disproportionation of chlorine in water? A. Cl2(g) + H2O(l) Cl-(aq) + H2(g) + ClO+(aq) B. Cl2(g) + H2O(l) Cl-(aq) + H+(aq) + HClO(aq) C. Cl2(g) + H2O(l) Cl-(aq) + HCl(aq) + OH-(aq) D. Cl2(g) + 2H2O(l) Cl-(aq) + 3H+(aq) + HClO2(aq) E. Cl2(g) + 4H2O(l) Cl-(aq) + 8H+(aq) + ClO4-(aq)
7. What is the molecular formula for chlorous acid? A. HCl B. HClO C. H2ClO D. HClO2 E. HClO3
8. Which of the following lists of nonmetals contains only elements that can be found in nature in their elemental forms? A. chlorine, nitrogen, oxygen B. silicon, arsenic, sulfur C. nitrogen, phosphorus, arsenic D. bromine, carbon, oxygen E. nitrogen, oxygen, sulfur
9. Fluorine and chlorine gases are primarily produced by A. distillation. B. oxidation by electrolysis. C. oxidation by hydronium ion. D. reduction by carbon monoxide. E. roasting with copper(II) sulfide.
10. Write a balanced net ionic equation for the reaction of chlorine gas with aqueous potassium iodide. A. Cl2(g) + 2KI(aq) ® 2KCl(s) + I2(aq) B. Cl2(g) + 2I-(aq) ® 2ICl(aq) C. Cl2(g) + 2I-(aq) ® 2Cl-(aq) + I2(aq) D. Cl2(g) + 2I-(aq) ® Cl2I(aq) + I-(aq) E. Cl2(g) + K+(aq) + I-(aq) ® KCl(s) + ICl(g)
11. Which of the following statements are true about the reaction below? NH3 + BF3 ® NH3BF3 1. 2. 3.
Ammonia acts as a Lewis base. Ammonia acts as a Brønsted-Lowry base. Ammonia acts as a reducing agent.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 2 and 3 12. Which of the following nitrogen containing compounds or ions is not capable of acting as an oxidizing agent? A. NO3B. NO2C. NH3 D. N2 E. NO2
13. Hydrazine is made commercially by reacting aqueous ammonia with hypochlorite ion. Write a balanced equation for this reaction. A. 2NH3(aq) + ClO4-(aq) ® N2H4(aq) + ClO3-(aq) + H2O(l) B. 2NH3(aq) + ClO3-(aq) ® N2H4(aq) + ClO2-(aq) + H2O(l) C. 2NH3(aq) + ClO2-(aq) ® N2H4(aq) + Cl-(aq) + H2O2(aq) D. 2NH3(aq) + ClO-(aq) ® N2H4(aq) + Cl-(aq) + H2O(l) E. 2NH3(aq) + ClO-(aq) ® N2H4(aq) + H2(aq) + ClO-(aq)
14. Write a balanced equation for the reaction of gaseous ammonia with hydrogen chloride. A. NH3(g) + HCl(g) ® NH2Cl(g) + H2(g) B. NH3(g) + HCl(g) ® NH4Cl(s) C. NH3(g) + 3HCl(g) ® NCl3(g) + 3H2(g) D. NH3(g) + 2HCl(g) ® NH5(g) + Cl2(g) E. 2NH3(g) + 2HCl(g) ® N2(g) + 4H2(g) + Cl2(g)
15. Which of the following statements best describe the role of hydrogen sulfide in the reaction below? 2H2S(aq) + O2(g) ® 2S(s) + 2H2O(l) 1. 2. 3.
H2S is a reducing agent. H2S is a Brønsted-Lowry base. H2S is a precipitating agent.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 2 and 3 16. Write a half-reaction for the reduction of aqueous hydrogen peroxide in an acidic solution. A. H2O2(aq) + 2e- ® H2(g) + O22-(aq) B. 2H2O2(aq) + 2e- ® O2(g) + 2H2O(l) C. H2O2(aq) + 2e- ® 2OH-(aq) D. H2O2(aq) + 2H+(aq) + 2e- ® 2H2O(l) E. Hydrogen peroxide cannot be reduced under acidic conditions.
17. Write a half-reaction for the oxidation of hydrogen peroxide. A. H2O2(aq) ® 2OH-(aq) + 2eB. H2O2(aq) ® H2(g) + O2(g) C. H2O2(aq) ® H2(g) + O2(g) + 2eD. 2H2O2(aq) ® O2(g) + 2H2O(l) + 2eE. H2O2(aq) ® O2(g) + 2H+(aq) + 2e-
18. What is the overall reaction for the disproportionation of hydrogen peroxide in water? A. H2O2(aq) ® 2OH-(aq) B. H2O2(aq) ® O2(g) + H2(g) C. 2H2O2(aq) ® O2(g) + 2H2O(l) D. H2O2(aq) ® O2(g) + 2H+(aq) E. H2O2(aq) + H2O(l) ® 4H+(aq) + O3(g)
19. Write a balanced chemical equation for the reaction of hydrogen peroxide and iodide ion in an acidic solution. A. H2O2(aq) + 2I-(aq) + 2H+(aq) ® I2(s) + 2H2O(l) B. H2O2(aq) + 2I-(aq) ® I2(s) + 2OH-(aq) C. H2O2(aq) + I2(s) ® 2I-(aq) + H2(g) + O2(g) D. H2O2(aq) + I2(s) + 2H+(aq) ® 2HI(s) + 2OH-(aq) E. H2O2(aq) + 2I-(aq) ® I2(s) + H2(g) + O2(g)
20. Write a balanced chemical equation for the reaction of aqueous solutions of hydrofluoric acid and potassium carbonate. A. 2HF(aq) + CO32-(aq) ® CO2(g) + H2O(l) + 2F-(aq) B. 2HF(aq) + K2CO3(aq) ® F2CO3(s) + H2(g) C. 2HF(aq) + K2CO3(s) ® 2KF(s) + H2CO3(s) D. 2H+(aq) + CO32-(aq) ® H2CO3(s) E. 2H+(aq) + CO32-(aq) ® CO2(g) + H2O(l)
21. Which of the following acids is never stored in glass containers because it reacts with silicon dioxide and ionic silicates? A. HNO3 B. HClO4 C. H2SO4 D. HF E. HI
22. All of the following oxidation states are observed for sulfur EXCEPT _____. A. -2 B. +2 C. +4 D. +6 E. +7
23. Which reddish-brown gas is a major factor in the formation of photochemical smog? A. NO B. NO2 C. SO2 D. CO2 E. CO
24. Which of the following molecules are paramagnetic: N2O, NO, and NO2? A. N2O only B. NO only C. NO2 only D. N2O and NO E. NO and NO2
25. Which of the following molecules has the ability to dilate blood vessels, thereby lowering blood pressure? A. NO B. N2O C. CO D. HCl E. SO3
26. Write a balanced chemical equation for the reaction of N2O5 and water. A. N2O5(s) + H2O(l) ® N2O4(s) + H2O2(l) B. N2O5(s) + H2O(l) ® N2O(g) + H2O(l) + 2O2(g) C. N2O5(s) + 3H2O(l) ® NH3(g) + 4O2(g) D. N2O5(s) + H2O(l) ® 2HNO3(l) E. N2O5(s) + H2O(l) ® 2HNO2(l) + O2(g)
27. Write a balanced chemical equation for the combustion of white phosphorus. A. P4(s) + 8O2(g) ® 4PO4(g) B. P4(s) + 3O2(g) ® 2P2O3(s) C. P4(s) + 5O2(g) ® P4O10(s) D. P4(s) + 8O2(g) ® 4PO43-(g) E. P4(s) + O2(g) ® 2P2O(g)
28. In which of the following compounds does phosphorus have an oxidation number of +3? A. P4S3 B. H3PO3 C. H3PO4 D. P4O10 E. PH3
29. Nonmetal oxides react with water to form A. solids. B. fullerenes. C. Lewis bases. D. Brønsted-Lowry acids. E. precipitates.
30. Write a balanced chemical equation for the reaction of dinitrogen trioxide with water. A. N2O3(g) + H2O(l) ® N2O4(g) + H2(g) B. N2O3(g) + H2O(l) ® 2HNO2(aq) C. N2O3(g) + 3H2O(l) ® 2HNO3(aq) + 2H2(g) D. N2O3(g) + H2O(l) ® N2H2(g) + 2O2(g) E. N2O3(g) + 3H2O(l) ® 2NH3(g) + 3O2(g)
31. Write a balanced chemical equation for the reaction of sulfur trioxide with water. A. SO3(g) + H2O(l) ® S(s) + H2(g) + 2O2(g) B. SO3(g) + H2O(l) ® H2SO3(aq) C. SO3(g) + H2O(l) ® SO2(g) + H2O2(g) D. SO3(g) + H2O(l) ® H2S(aq) + 2O2(g) E. SO3(g) + H2O(l) ® H2SO4(aq)
32. Which of the following oxoacids has the largest Ka value? A. HClO4 B. HClO3 C. HClO2 D. HClO E. All of the above oxoacids have similar Ka values.
33. Place the following acids in order from smallest to largest Ka value: HClO, HBrO, and HIO. A. HClO < HBrO < HIO B. HClO < HIO < HBrO C. HBrO < HClO < HIO D. HBrO < HIO < HClO E. HIO < HBrO < HClO
34. Calculate the pH of a 0.35 M solution of sodium hypobromite, NaClO. The acid dissociation constant for HBrO is 2.6 ´ 10-9. A. 2.94 B. 4.96 C. 8.13 D. 9.04 E. 11.06
35. H3PO4 is a weak, triprotic acid. The acid dissociation constants are Ka1 = 7.1 ´ 10-3, Ka2 = 6.2 ´ 10-8, and Ka3 = 4.5 ´ 10-13. Calculate the equilibrium constant for the reaction below. H3PO4(aq) PO43-(aq) + 3H+(aq) A. 2.0 ´ 10-22 B. 4.4 ´ 10-10 C. 2.0 ´ 10-8 D. 8.7 ´ 10-6 E. 7.1 ´ 10-3
36. Which of the following oxoanions of chlorine cannot act as an oxidizing agent? A. HClO B. HClO2 C. HClO3 D. HClO4 E. All of the above can act as oxidizing agents.
37. Which of the following nitrogen containing compounds is not able to act as an oxidizing agent? A. HNO3 B. NH3 C. HNO2 D. N2H4 E. All of the above can act as oxidizing agents.
38. Write a balanced chemical equation for the disproportionation of chlorate ion in water. The products are chloride and perchlorate ions. A. 2ClO3-(aq) + H2O(l) ® Cl-(aq) + ClO4-(aq) + 2H+(aq) B. 3ClO3-(aq) + 2H2O(l) ® Cl-(aq) + 2ClO4-(aq) + 4H+(aq) C. 4ClO3-(aq) ® Cl-(aq) + 3ClO4-(aq) D. 4ClO3-(aq) ® 2ClO2-(aq) + 2ClO4-(aq) E. 5ClO3-(aq) + 6H+(aq) ® 2Cl-(aq) + 3ClO4-(aq) + 3H2O(l)
39. The standard reduction potential of H2O2 = +1.763 V. Which of the following species will be oxidized by hydrogen peroxide?
Co2+(aq) ® Co3+(aq) + 2e Ag(s) ® Ag+(aq) + eCo(s) ® Co2+(aq) + 2eZn(s) ® Zn2+(aq) + 2eNa(s) ® Na+(aq) + e-
+1.953 V +0.799 V -0.282 V -0.762 V -2.714 V
A. Co2+ B. Co2+ and Ag(s) C. Co(s) and Co2+ D. Zn(s) and Na(s) E. Ag(s), Co(s), Zn(s), and Na(s) 40. The first step in the synthesis of nitric acid involves the reaction of ammonia with oxygen to produce nitrogen monoxide. 4NH3(g) + 5O2(g) ® 4NO(g) + 6H2O(g) What volume of NO is produced if 14.0 L NH3 is reacted with 15.0 L O2? Assume the reaction is 100% efficient. A. 12.0 L B. 14.0 L C. 15.0 L D. 23.2 L E. 29.0 L
41. Write a balanced chemical equation for the oxidation of iron metal with nitric acid. Assume the reaction products are Fe3+ and NO gas. A. Fe(s) + NO3-(aq) ® Fe3+(aq) + NO(g) + O2(g) B. Fe(s) + HNO3(aq) ® Fe3+(aq) + NO(g) + OH-(aq) C. Fe(s) + 4H+(aq) + NO3-(aq) ® Fe3+(aq) + NO(g) + 2H2O(l) D. Fe(s) + H+(aq) + NO3-(aq) ® Fe3+(aq) + NO(g) + O2H-(aq) E. 2Fe(s) + 12H+(aq) + 3NO3-(aq) ® 2Fe3+(aq) + 3NO(g) + 6H2O(l)
42. Write a balanced chemical equation for the oxidation of chromium metal with nitric acid. Assume the reaction products are Cr3+ and NH4+. A. H+(aq) + NO3-(aq) + Cr(s) ® NH4+(aq) + Cr3+(aq) B. 4H+(aq) + NO3-(aq) + Cr(s) ® NH4+(aq) + Cr3+(aq) C. 10H+(aq) + NO3-(aq) + 3Cr(s) ® NH4+(aq) + 3Cr3+(aq) + 3H2O(l) D. 10H+(aq) + NO3-(aq) + Cr(s) ® NH4+(aq) + Cr3+(aq) + 3H2O(l) E. 30H+(aq) + 3NO3-(aq) + 8Cr(s) ® 3NH4+(aq) + 8Cr3+(aq) + 9H2O(l)
43. Which of the following compounds will react with water to produce a basic solution: NH3, H2O2, and H2S? A. NH3 only B. H2O2 only C. H2S only D. NH3 and H2S E. H2O2 and H2S
44. All of the following statements concerning oxoanions and oxoacids are true EXCEPT A. MnO4- can never act as an oxidizing agent. B. Cr2O72- can never act as a reducing agent. C. ClO4- is a stronger oxidizing agent than ClO-. D. ClO2- can behave as either an oxidizing or a reducing agent. E. HClO3 is a stronger acid than HClO2.
45. Nitric acid solutions slowly turn brown when exposed to sunlight. Which of the following reactions is responsible? A. NO3-(aq) ® NO(g) + O2(g) B. 2NO3-(aq) ® 2NO2(g) + O2(g) C. 2H+(aq) + NO3-(aq) ® NO2(g) + H2O(l) D. 4H+(aq) + 4NO3-(aq) ® 4NO2(g) + 2H2O(l) + O2(g) E. 4H+(aq) + NO3-(aq) ® NO(g) + 2H2(g) + O2(g)
46. Hydrazine, N2H4, is produced by the Raschig process-the oxidation of ammonia with alkaline ____. A. HNO3(aq) B. NaOCl(aq) C. N2(g) D. HPO42–(aq) E. N2O(g)
47. Sulfuric acid can act as 1. 2. 3.
a drying agent. an oxidizing agent. a Brønsted-Lowry acid.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 48. Write a balanced equation for the reaction of table sugar, C12H22O11, with concentrated sulfuric acid. A. C12H22O11(s) ® 12C(s) + 11H2O(l) B. C12H22O11(s) + H2SO4(aq) ® C12H23O11+(s) + HSO4-(aq) C. C12H22O11(s) + H2SO4(aq) + 6O2(g) ® 12CO2(g) + SO2(g) + H2O(l) D. C12H22O11(s) + 2H+(aq) + SO42-(aq) ® C12H24O112+(s) + SO42-(aq) E. C12H22O11(s) + 2H+(aq) + SO42-(aq) ® 12CO(g) + SO2(g) + 2OH+(aq) + H2O(l)
49. Which of the following statements is INCORRECT? A. Sodium dihydrogen phosphate, NaH2PO4, is used in acidic cleaners. B. Sodium phosphate, Na3PO4, is used in strongly basic cleaners. C. Calcium dihydrogen phosphate, Ca(H2PO4)2 is used in fertilizers. D. The principal use of Na2HPO4 is as an emulsifier in the manufacture of cheese. E. Phosphoric acid is a strong acid, a powerful oxidizing agent, and a drying agent.
50. Write a balanced equation for the reaction of concentrated sulfuric acid with copper metal. Assume two products are sulfur dioxide and Cu2+. A. Cu(s) + SO42-(aq) ® Cu2+(aq) + O2(g) + SO2(g) B. Cu(s) + 4H+(aq) + SO42-(aq) ® CuS(s) + 2O2(g) C. Cu(s) + 2H+(aq) + SO42-(aq) ® Cu2+(aq) + H2O(l) + SO2(g) D. Cu(s) + 4H+(aq) + SO42-(aq) ® Cu2+(aq) + 2H2O(l) + SO2(g) E. Cu(s) + 8H+(aq) + SO42-(aq) ® Cu2+(aq) + 4H2O(l) + S(g)
51. Write a balanced equation for the reaction of Al3+ with aqueous ammonia. A. Al3+(aq) + 2NH3(aq) ® Al(s) + 3H2(g) + N2(g) B. Al3+(aq) + 3NH3(aq) + 3H2O(l) ® Al(OH)3(s) + 3NH4+(aq) C. Al3+(aq) + 2NH3(aq) ® Al(NH3)23+(aq) D. Al3+(aq) + 3NH3(aq) + 3H+(aq) ® Al(NH4)36+(aq) E. Al3+(aq) + 6NH3(aq) ® Al(NH3)63+(aq)
52. Write a balanced equation for the reaction of Ni2+ with H2S. A. Ni2+(aq) + 2H2S(aq) ® Ni(HS)2(s) + 2H+(aq) B. Ni2+(aq) + H2S(aq) ® NiH2(s) + S2-(aq) C. Ni2+(aq) + H2S(aq) ® NiS(s) + 2H+(aq) D. Ni2+(aq) + H2S(aq) ® Ni(s) + S(s) + 2H+(aq) E. No reaction will occur.
53. Hydrogen sulfide is used to remove sulfur dioxide in power plant smokestacks. The reaction produces elemental sulfur and water. Write a balanced chemical equation for this reaction. A. H2S(l) + SO2(g) ® 2HSO(s) B. H2S(l) + SO2(g) ® 2S(s) + 2OH-(aq) C. 2H2S(l) + SO2(g) ® 3S(s) + 2H2O(l) D. H2S(l) + SO2(g) ® 2S(s) + H2O2(aq) E. H2S(l) + SO2(g) + 2H+(aq) ® 2S(s) + 2H2O(l)
54. In modern forensic laboratories, arsenic poisoning is detected by testing A. a strand of hair. B. samples of brain tissue. C. samples of liver tissue. D. heart tissue. E. bone fragments.
55. As one might predict, one allotrope of arsenic has the formula A. As4 B. As8 C. As2 D. As3 E. As10
56. The average concentration of bromine (as bromide) in seawater is 65 ppm. Calculate the volume of seawater (d = 64.0 lb/ft3) in cubic feet required to produce one kilogram of liquid bromine. A. 109 ft3 B. 448 ´ 103 ft3 C. 240 ft3 D. 154 ´ 102 ft3 E. 529 ft3
Chapter 22--Organic Chemistry 1. Saturated hydrocarbons are also known as A. methanes. B. alkanes. C. alkynes. D. alkenes. E. ketones.
2. What is the general formula for an alkane? A. CnH2n+2 B. CnH2n C. CnHn–2 D. Cn+2Hn E. CnH2n–2
3. What are the approximate bond angles between the carbon and hydrogen atoms in C3H8? A. 60 B. 90 C. 109.5 D. 120 E. 90 and 120
4. What is the hybridization of the carbon atoms in an alkane? A. s B. p C. sp D. sp2 E. sp3
5. Which of the statements concerning the two molecules below are correct?
1. 2. 3.
The molecules are unsaturated hydrocarbons. The molecules are structural isomers. The molecules are alkanes.
A. 1 only B. 2 only C. 3 only D. 2 and 3 E. 1, 2, and 3 6. Which of the following statements concerning structural isomers is/are correct? 1. 2. 3.
Structural isomers have the same elemental composition, but the atoms are linked in different ways. Structural isomers have identical physical properties, but different chemical properties. Structural isomers have identical structures, but contain different isotopes of the same elements.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 7. How many structural isomers exist for C5H12? A. 2 B. 3 C. 4 D. 5 E. 6
8. What type of intermolecular forces or bonds dominate in alkanes? A. dispersion forces B. dipole forces C. hydrogen bonds D. covalent bonds E. ionic bonds
9. Cylinders of bottled gas used in camp stoves and grills contain liquid A. methane. B. ethane. C. propane D. acetylene. E. ethene.
10. Which of the following alkanes will have the lowest boiling point?
A.
B.
C.
D. E. All the above alkanes have the same boiling point.
11. What process is used to separate petroleum into various products, such as gasoline, kerosene, and fuel oil? A. mass spectroscopy B. filtration C. gas chromatography D. column chromatography E. fractional distillation
12. Gasoline is A. 90% isooctane and 10% antiknocking compounds. B. primarily composed of methyl-t-butyl ether (MTBE). C. primarily composed of isooctane. D. a mixture of hydrocarbons in the C5 to C12 range. E. more volatile than propane, but less volatile than butane.
13. The octane number of a gasoline is a measure of A. the mass of octane in a 100 gram sample of gasoline. B. its resistance to knock (i.e., premature ignition). C. the mass of a gasoline sample divided by the molar mass of octane. D. the vapor pressure of gasoline at 298 K. E. the percentage of octane in a sample of gasoline.
14. Which of the following alkanes will have the highest boiling point?
A.
B.
C.
D. E. Molecules a, b, and c have equally high boiling points.
15. What is the name of the following compound?
A. 3,4-diethyl-5-methylheptane B. 3,4,5-heptane C. 3-ethyl-5-(2-butyl)hexane D. 3-(2-butyl)-3-ethylhexane E. 3,5,6-trimethylheptane
16. What is the IUPAC name of the following compound?
A. 2,4-dimethylhexane B. 2-ethyl-4-methylpentane C. 2,4-methylhexane D. 3,5-dimethylhexane E. 2,4-octane
17. What is the name of the following compound?
A. pentanol B. cyclopentane C. cyclohexane D. cyclopentene E. 1,5-pentane
18. What is the name of the following compound?
A. trans-acetylene B. cis-1,3-pentadiene C. trans-2,4-pentadiene D. cis-4-methyl-1,3-butadiene E. cis-4-methyl-1,3-butene
19. Which of the following (non-cyclic) molecules is an alkyne? A. CH4 B. C2H2 C. C2H6 D. C3H6 E. C6H14
20. How many p bonds are present in (the noncyclic hydrocarbon) C3H4? A. 0 B. 1 C. 2 D. 3 E. 6
21. What is the hybridization of each carbon atom in acetylene, C2H2? A. sp B. sp2 C. sp3 D. sp3d E. 2s and 2p
22. All of the following statements concerning ethene are true EXCEPT A. ethene is the most produced organic compound in the United States. B. ethene is the starting material for polyethylene. C. small quantities of ethene are used to ripen fruit. D. ethene is a chiral molecule. E. ethene has the molecular formula C2H4.
23. What is the IUPAC name of the following compound?
A. 3-methyl-1-butyne B. 3-methyl-1-butene C. 2-methyl-3-butene D. 2-methyl-4-butene E. 3,3-dimethyl-1-propene
24. What is the IUPAC name of the following compound?
A. 4-methyl-2-pentyne B. 4-methyl-2,3-pentyne C. 4-methyl-2,3-dipentyne D. 2-methyl-3-pentyne E. 2-methyl-3,4-pentyne
25. Which one of the following hydrocarbons is aromatic?
A.
B.
C.
D. E.
26. Which of the statements concerning benzene are true? 1. 2. 3.
Each carbon forms three sigma bonds. The p bonds are delocalized over the entire molecule. The hybridization of the carbon atoms is sp2.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 27. What is the hybridization of each carbon atom in benzene, C6H6? A. 2s B. 2p C. sp D. sp2 E. sp3
28. What is the name of the following benzene derivative?
A. 1-chlorobenzoic acid B. 5-chloroanaline C. 1-acetate-1-chlorobenzene D. 5-chlorobenzoic acid E. 1,3-chlorocarboxylic benzene
29. Which of the structures below has the common name o-dichlorobenzene (where o- is ortho)?
A.
B.
C.
D.
E.
30. What is the name of the following benzene derivative?
A. dibenzene B. 1,2-dibenzene C. toluene D. naphthalene E. aniline
31. Classify the following molecule according to its functional group.
A. alcohol B. aldehyde C. carboxylic acid D. ester E. ether
32. Classify the following molecule according to its functional group.
A. alcohol B. carboxylic acid C. ketone D. ester E. ether
33. Formulas for derivatives of hydrocarbons may be written as R-X, where R is a hydrocarbon lacking a hydrogen atom and X is a functional group. Which of the following formulas represents a ketone? A. ROH B. ROR' C. RCOR' D. RCO2R' E. RCHO
34. Classify the following molecule according to its functional group.
A. alcohol B. aldehyde C. alkane D. carboxylic acid E. ester
35. Classify the following molecule according to its functional group.
A. alcohol B. aldehyde C. carboxylic acid D. ether E. ketone
36. Formulas for derivatives of hydrocarbons may be written as R-X, where R is a hydrocarbon lacking a hydrogen atom and X is a functional group. Which of the following formulas represents an ester? A. ROH B. ROR' C. RCOR' D. RCO2R' E. RCHO
37. Which compounds often have strong, pleasant odors? A. alcohols B. aldehydes C. amines D. esters E. ethers
38. Which functional group does not contain a double bond to an oxygen atom? A. ester B. aldehyde C. ether D. amide E. ketone
39. How many alcohols have the chemical formula C4H10O? A. 1 B. 2 C. 3 D. 4 E. 5
40. How many ketones have the chemical formula C5H10O? A. 2 B. 3 C. 4 D. 5 E. 6
41. How many ethers have the chemical formula C5H12O? A. 3 B. 4 C. 5 D. 6 E. 7
42. How many different (non-cyclic) structural isomers exist for C6H12? A. 4 B. 8 C. 10 D. 12 E. 13
43. How many isomers exist for the following benzene derivative, C6H4ClBr? A. 2 B. 3 C. 4 D. 5 E. none
44. How many different structures (including structural and geometric isomers) exist for C2H2Cl2? A. 2 B. 3 C. 4 D. 5 E. 6
45. Which of the following molecules can form cis and trans isomers? 1. 2. 3.
CH3ClC=CH2 CH3ClC=CCH3Cl BrClC=CClCH3
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 2 and 3 46. For which one of the following molecules do geometric isomers exist? A. H2ClC-CHCl2 B. BrHC=CHBr C. ClCºCH D. H3C-CH2Cl E. H2C=CCl2
47. Optical isomerism occurs when A. a molecule's mirror image is superimposable. B. both cis and trans isomers exist in equal concentrations. C. two or more structural isomers of a molecule exist. D. at least one carbon atom in a molecule is bonded to four different atoms or groups. E. both enantiomers are present in a racemic mixture.
48. Which one of the following statements concerning isomers is INCORRECT? A. Pairs of nonsuperimposable molecules are called enantiomers. B. Enantiomers have identical melting and boiling points. C. Molecules with two or more geometric isomers are termed chiral pairs. D. Optical isomers rotate polarized light in opposite directions. E. Structural isomers have the same composition, but the atoms are linked in different ways.
49. Which of the following molecules has at least one chiral center?
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 2 and 3
50. Which of the following molecules have chiral centers?
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1 and 3
51. Write a balanced chemical equation for the reaction of ethene with oxygen. A. CH4(g) + 2O2(g) ® CO2(g) + 2H2O(g) B. 2C2H6(g) + 7O2(g) ® 4CO2(g) + 6H2O(g) C. 2C2H2(g) + 5O2(g) ® 4CO2(g) + 2H2O(g) D. C2H4(g) + 3O2(g) ® 2CO2(g) + 2H2O(g) E. C2H4(g) + 2O2(g) ® 2CO2(g) + 2H2(g)
52. What is the product of the addition of HCl to ethylene? A. chloroethane B. chloroethylene C. 1,2-dichloroethane D. 1,1-dichloroethane E. 1,1,2,2-tetrachloroethylene
53. Which of the following is an elimination reaction? A. C2H2(g) + H2(g) ® C2H4(g) B. CH3CH2CH2OH(l) ® CH3CHCH2(g) + H2O(l) C. C3H6(g) + HCl(g) ® C3H7Cl(l) D. CH3OH(aq) + HCOOH(aq) ® C2H5-O-CHO(l) + H2O(l) E. 2C2H2(g) + 5O2(g) ® 4CO2(g) + 2H2O(g)
54. Which of the following is a substitution reaction? A. C2H2(g) + H2(g) ® C2H4(g) B. C6H6(l) + HNO3(l) ® C6H5NO2(l) + H2O(l) C. CH3CH2CH2OH(l) ® CH3CHCH2(g) + H2O(l) D. CH3OH(aq) + HCOOH(aq) ® C2H5-O-CHO(l) + H2O(l) E. 2C2H2(g) + 5O2(g) ® 4CO2(g) + 2H2O(g)
55. Alkenes undergo addition reactions with halogens. What is the product of the reaction of propylene, C3H6, with Br2?
A.
B.
C.
D.
E.
56. The Cholesterol molecule is found in two types of complexes: low-density lipoproteins (LDL) and high-density lipoproteins (HDL). Which of these is considered "bad"? A. (LDL) B. (HDL) C. both (HDL) and (LDL), because they produce trans fat. D. Neither, only omega-3 fatty acids are bad E. both (HDL) and (LDL), because they form saturated fat.
Chapter 22--Organic Chemistry Key
1. Saturated hydrocarbons are also known as A. methanes. B. alkanes. C. alkynes. D. alkenes. E. ketones.
2. What is the general formula for an alkane? A. CnH2n+2 B. CnH2n C. CnHn–2 D. Cn+2Hn E. CnH2n–2
3. What are the approximate bond angles between the carbon and hydrogen atoms in C3H8? A. 60 B. 90 C. 109.5 D. 120 E. 90 and 120
4. What is the hybridization of the carbon atoms in an alkane? A. s B. p C. sp D. sp2 E. sp3
5. Which of the statements concerning the two molecules below are correct?
1. 2. 3.
The molecules are unsaturated hydrocarbons. The molecules are structural isomers. The molecules are alkanes.
A. 1 only B. 2 only C. 3 only D. 2 and 3 E. 1, 2, and 3 6. Which of the following statements concerning structural isomers is/are correct? 1. 2. 3.
Structural isomers have the same elemental composition, but the atoms are linked in different ways. Structural isomers have identical physical properties, but different chemical properties. Structural isomers have identical structures, but contain different isotopes of the same elements.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 7. How many structural isomers exist for C5H12? A. 2 B. 3 C. 4 D. 5 E. 6
8. What type of intermolecular forces or bonds dominate in alkanes? A. dispersion forces B. dipole forces C. hydrogen bonds D. covalent bonds E. ionic bonds
9. Cylinders of bottled gas used in camp stoves and grills contain liquid A. methane. B. ethane. C. propane D. acetylene. E. ethene.
10. Which of the following alkanes will have the lowest boiling point?
A.
B.
C.
D. E. All the above alkanes have the same boiling point.
11. What process is used to separate petroleum into various products, such as gasoline, kerosene, and fuel oil? A. mass spectroscopy B. filtration C. gas chromatography D. column chromatography E. fractional distillation
12. Gasoline is A. 90% isooctane and 10% antiknocking compounds. B. primarily composed of methyl-t-butyl ether (MTBE). C. primarily composed of isooctane. D. a mixture of hydrocarbons in the C5 to C12 range. E. more volatile than propane, but less volatile than butane.
13. The octane number of a gasoline is a measure of A. the mass of octane in a 100 gram sample of gasoline. B. its resistance to knock (i.e., premature ignition). C. the mass of a gasoline sample divided by the molar mass of octane. D. the vapor pressure of gasoline at 298 K. E. the percentage of octane in a sample of gasoline.
14. Which of the following alkanes will have the highest boiling point?
A.
B.
C.
D. E. Molecules a, b, and c have equally high boiling points.
15. What is the name of the following compound?
A. 3,4-diethyl-5-methylheptane B. 3,4,5-heptane C. 3-ethyl-5-(2-butyl)hexane D. 3-(2-butyl)-3-ethylhexane E. 3,5,6-trimethylheptane
16. What is the IUPAC name of the following compound?
A. 2,4-dimethylhexane B. 2-ethyl-4-methylpentane C. 2,4-methylhexane D. 3,5-dimethylhexane E. 2,4-octane
17. What is the name of the following compound?
A. pentanol B. cyclopentane C. cyclohexane D. cyclopentene E. 1,5-pentane
18. What is the name of the following compound?
A. trans-acetylene B. cis-1,3-pentadiene C. trans-2,4-pentadiene D. cis-4-methyl-1,3-butadiene E. cis-4-methyl-1,3-butene
19. Which of the following (non-cyclic) molecules is an alkyne? A. CH4 B. C2H2 C. C2H6 D. C3H6 E. C6H14
20. How many p bonds are present in (the noncyclic hydrocarbon) C3H4? A. 0 B. 1 C. 2 D. 3 E. 6
21. What is the hybridization of each carbon atom in acetylene, C2H2? A. sp B. sp2 C. sp3 D. sp3d E. 2s and 2p
22. All of the following statements concerning ethene are true EXCEPT A. ethene is the most produced organic compound in the United States. B. ethene is the starting material for polyethylene. C. small quantities of ethene are used to ripen fruit. D. ethene is a chiral molecule. E. ethene has the molecular formula C2H4.
23. What is the IUPAC name of the following compound?
A. 3-methyl-1-butyne B. 3-methyl-1-butene C. 2-methyl-3-butene D. 2-methyl-4-butene E. 3,3-dimethyl-1-propene
24. What is the IUPAC name of the following compound?
A. 4-methyl-2-pentyne B. 4-methyl-2,3-pentyne C. 4-methyl-2,3-dipentyne D. 2-methyl-3-pentyne E. 2-methyl-3,4-pentyne
25. Which one of the following hydrocarbons is aromatic?
A.
B.
C.
D. E.
26. Which of the statements concerning benzene are true? 1. 2. 3.
Each carbon forms three sigma bonds. The p bonds are delocalized over the entire molecule. The hybridization of the carbon atoms is sp2.
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 1, 2, and 3 27. What is the hybridization of each carbon atom in benzene, C6H6? A. 2s B. 2p C. sp D. sp2 E. sp3
28. What is the name of the following benzene derivative?
A. 1-chlorobenzoic acid B. 5-chloroanaline C. 1-acetate-1-chlorobenzene D. 5-chlorobenzoic acid E. 1,3-chlorocarboxylic benzene
29. Which of the structures below has the common name o-dichlorobenzene (where o- is ortho)?
A.
B.
C.
D.
E.
30. What is the name of the following benzene derivative?
A. dibenzene B. 1,2-dibenzene C. toluene D. naphthalene E. aniline
31. Classify the following molecule according to its functional group.
A. alcohol B. aldehyde C. carboxylic acid D. ester E. ether
32. Classify the following molecule according to its functional group.
A. alcohol B. carboxylic acid C. ketone D. ester E. ether
33. Formulas for derivatives of hydrocarbons may be written as R-X, where R is a hydrocarbon lacking a hydrogen atom and X is a functional group. Which of the following formulas represents a ketone? A. ROH B. ROR' C. RCOR' D. RCO2R' E. RCHO
34. Classify the following molecule according to its functional group.
A. alcohol B. aldehyde C. alkane D. carboxylic acid E. ester
35. Classify the following molecule according to its functional group.
A. alcohol B. aldehyde C. carboxylic acid D. ether E. ketone
36. Formulas for derivatives of hydrocarbons may be written as R-X, where R is a hydrocarbon lacking a hydrogen atom and X is a functional group. Which of the following formulas represents an ester? A. ROH B. ROR' C. RCOR' D. RCO2R' E. RCHO
37. Which compounds often have strong, pleasant odors? A. alcohols B. aldehydes C. amines D. esters E. ethers
38. Which functional group does not contain a double bond to an oxygen atom? A. ester B. aldehyde C. ether D. amide E. ketone
39. How many alcohols have the chemical formula C4H10O? A. 1 B. 2 C. 3 D. 4 E. 5
40. How many ketones have the chemical formula C5H10O? A. 2 B. 3 C. 4 D. 5 E. 6
41. How many ethers have the chemical formula C5H12O? A. 3 B. 4 C. 5 D. 6 E. 7
42. How many different (non-cyclic) structural isomers exist for C6H12? A. 4 B. 8 C. 10 D. 12 E. 13
43. How many isomers exist for the following benzene derivative, C6H4ClBr? A. 2 B. 3 C. 4 D. 5 E. none
44. How many different structures (including structural and geometric isomers) exist for C2H2Cl2? A. 2 B. 3 C. 4 D. 5 E. 6
45. Which of the following molecules can form cis and trans isomers? 1. 2. 3.
CH3ClC=CH2 CH3ClC=CCH3Cl BrClC=CClCH3
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 2 and 3 46. For which one of the following molecules do geometric isomers exist? A. H2ClC-CHCl2 B. BrHC=CHBr C. ClCºCH D. H3C-CH2Cl E. H2C=CCl2
47. Optical isomerism occurs when A. a molecule's mirror image is superimposable. B. both cis and trans isomers exist in equal concentrations. C. two or more structural isomers of a molecule exist. D. at least one carbon atom in a molecule is bonded to four different atoms or groups. E. both enantiomers are present in a racemic mixture.
48. Which one of the following statements concerning isomers is INCORRECT? A. Pairs of nonsuperimposable molecules are called enantiomers. B. Enantiomers have identical melting and boiling points. C. Molecules with two or more geometric isomers are termed chiral pairs. D. Optical isomers rotate polarized light in opposite directions. E. Structural isomers have the same composition, but the atoms are linked in different ways.
49. Which of the following molecules has at least one chiral center?
A. 1 only B. 2 only C. 3 only D. 1 and 3 E. 2 and 3
50. Which of the following molecules have chiral centers?
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1 and 3
51. Write a balanced chemical equation for the reaction of ethene with oxygen. A. CH4(g) + 2O2(g) ® CO2(g) + 2H2O(g) B. 2C2H6(g) + 7O2(g) ® 4CO2(g) + 6H2O(g) C. 2C2H2(g) + 5O2(g) ® 4CO2(g) + 2H2O(g) D. C2H4(g) + 3O2(g) ® 2CO2(g) + 2H2O(g) E. C2H4(g) + 2O2(g) ® 2CO2(g) + 2H2(g)
52. What is the product of the addition of HCl to ethylene? A. chloroethane B. chloroethylene C. 1,2-dichloroethane D. 1,1-dichloroethane E. 1,1,2,2-tetrachloroethylene
53. Which of the following is an elimination reaction? A. C2H2(g) + H2(g) ® C2H4(g) B. CH3CH2CH2OH(l) ® CH3CHCH2(g) + H2O(l) C. C3H6(g) + HCl(g) ® C3H7Cl(l) D. CH3OH(aq) + HCOOH(aq) ® C2H5-O-CHO(l) + H2O(l) E. 2C2H2(g) + 5O2(g) ® 4CO2(g) + 2H2O(g)
54. Which of the following is a substitution reaction? A. C2H2(g) + H2(g) ® C2H4(g) B. C6H6(l) + HNO3(l) ® C6H5NO2(l) + H2O(l) C. CH3CH2CH2OH(l) ® CH3CHCH2(g) + H2O(l) D. CH3OH(aq) + HCOOH(aq) ® C2H5-O-CHO(l) + H2O(l) E. 2C2H2(g) + 5O2(g) ® 4CO2(g) + 2H2O(g)
55. Alkenes undergo addition reactions with halogens. What is the product of the reaction of propylene, C3H6, with Br2?
A.
B.
C.
D.
E.
56. The Cholesterol molecule is found in two types of complexes: low-density lipoproteins (LDL) and high-density lipoproteins (HDL). Which of these is considered "bad"? A. (LDL) B. (HDL) C. both (HDL) and (LDL), because they produce trans fat. D. Neither, only omega-3 fatty acids are bad E. both (HDL) and (LDL), because they form saturated fat.
Chapter 23--Organic Polymers, Natural and Synthetic 1. All of the following statements concerning polyethylene are correct EXCEPT A. plastic bags at supermarkets are made from polyethylene. B. linear polyethylene is more rigid than branched polyethylene. C. linear polyethylene is referred to as high-density polyethylene. D. polyethylene is made by addition reaction of ethylene monomer units. E. the branches in branched polyethylene are primarily carboxylic acid groups.
2. Low density polyethylene (LDPE) is most likely to be used in A. sandwich bags. B. bottle caps. C. rigid pipes. D. contact lenses. E. fabrics and apparel.
3. PVC is a stiff, rugged polymer made from the addition reaction of ____. A. C2H4 B. C3H6 C. CH2CHCl D. CH2CHCN E. C2F4
4. Teflon is made from the addition reaction of ____. A. C2H4 B. C3H6 C. CH2CHCl D. CH2CHCN E. C2F4
5. Polyacrylonitrile is an addition polymer made from the acrylonitrile monomer, CH2CHCN. How many monomer units are in a polyacrylonitrile polymer that has a molar mass of 8.33 ´ 104 g/mol? A. 7.90 ´ 102 B. 1.57 ´ 103 C. 3.14 ´ 103 D. 8.33 ´ 104 E. 4.42 ´ 106
6. Which one of the following polymers is made by condensation reaction? A. Teflon B. poly(ethylene terephthalate) C. polyethylene D. polystyrene E. polyvinyl chloride
7. In principle, the acrylonitrile monomer (CH2CHCN) can polymerize in which of the following ways? 1. 2. 3.
A head-to-tail polymer A head-to-head, tail-to-tail polymer A random polymer
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 8. The reaction between an amine and a carboxylic acid is a condensation. What small molecule is formed as a byproduct of this reaction? A. H2O B. H2 C. CO2 D. O2 E. CO
9. When a diamine reacts with a dicarboxylic acid, a ____ is formed? A. polyamide B. polyester C. disaccharide D. carbohydrate E. saccharide
10. The polymerization of 1,2-dichloroethylene, CHClCHCl, will yield which of the following polymers?
A.
B.
C.
D.
E.
11. Which of the following molecules will react with oxalic acid to produce a polymer with the following repeating unit?
A.
B.
C.
D.
E.
12. What is the monomer used to make the following polymer?
A.
B.
C.
D.
E.
13. Which of the following statements is/are CORRECT concerning Nylon-66? 1. 2. 3.
The polymer forms at the interface of two insoluble liquds. One monomer is a dicarboxylic acid, the other monomer is a diamine. The polymerization reaction is a condensation with ammonia formed as a byproduct.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3
14. Glycine, H2NCH2COOH, can polymerize via a condensation reaction. What simple molecule is a byproduct of this polymerization reaction? A. H2 B. NH3 C. N2 D. H2O E. O2
15. The general formula of a carbohydrate is ____. A. CnH2nOn B. CnH2n+2On C. CnH2nO2n D. Cn(H2O)m E. Cn(OH)m
16. Which of the following statements are correct? 1. 2. 3.
Monosaccharides contain either five or six carbon atoms per molecule. Disaccharides are formed when two monosaccharide units combine with the elimination of water. Polysaccharides generally contain anywhere from several hundred to several thousand monosaccharide units.
A. 1 only B. 2 only C. 3 only D. 2 and 3 E. 1, 2, and 3 17. Which of the following statements concerning glucose are correct? 1. 2. 3.
As a pure solid, glucose exists primarily as six-membered ring molecules. In the six-membered ring structure, a CH2OH group is bonded to carbon atom 5. a-glucose and b-glucose are not enantiomers. Although they have several chiral carbon atoms, they differ in configuration at only one carbon atom.
A. 1 only B. 2 only C. 3 only D. 2 and 3 E. 1, 2, and 3 18. Which of the following statements concerning glucose are correct? 1. 2. 3.
a-Glucose and b-glucose each have one chiral carbon atom. In aqueous solution, a-glucose and b-glucose exist as an equilibrium mixture. a-Glucose and b-glucose are enantiomers.
A. 1 only B. 2 only C. 3 only D. 2 and 3 E. 1, 2, and 3 19. Sucrose is commonly known as ____. A. aspartame B. Splenda C. NutraSweet D. saccharine E. sugar
20. Fructose is A. commonly known as blood sugar. B. a monosaccharide found in honey and fruit juices. C. a disaccharide formed from two glucose molecules. D. a polysaccharide formed from several hundred glucose molecules. E. an enzyme that catalyzes the decomposition of glucose.
21. Which of the following is a polysaccharide? A. a-glucose B. fructose C. cellulose D. sucrose E. a-maltose
22. Which of the following statements concerning starch is INCORRECT? A. Starch is a polysaccharide found in plants. B. Starch is abundant in corn and potatoes. C. Amylose comprises about 20% of starch and it is insoluble in water. D. Amylopectin, a highly branched component of starch, is soluble in water. E. Amylose is composed of a-glucose polymers; amylopectin is composed of b-glucose polymers.
23. Which of the following statements concerning cellulose is INCORRECT? A. Humans lack the enzymes needed to catalyze the hydrolysis of cellulose. B. Cellulose contains unbranched chains of roughly 50 - 300 glucose molecules. C. The molecular structure of cellulose allows for hydrogen bonding between polymer chains. D. The major industrial source of cellulose is wood. E. In cellulose, the oxygen bridge between glucose units is in the beta position.
24. Which amino acid contains no chiral carbon atoms? A. proline B. glutamine C. lysine D. glycine E. leucine
25. How many of the 20 common a-amino acids are considered "essential" (i.e., required by humans for protein construction). A. 2 B. 5 C. 10 D. 13 E. 17
26. For an amino acid, proton transfer occurs from an carboxylic acid group to its amine group. The resulting species, which has both a +1 and a -1 charge, is called a(n) _____ . A. amide B. zwitterion C. cation D. anion E. protein
27. Using the acid dissociation constants for glycine (Ka1 = 4.6 ´ 10-3 and Ka2 = 2.5 ´ 10-10), calculate the [Z]/[C+] ratio at pH 1.00. (Z = Zwitterion, C+ = fully protonated cation) A. 2.5 ´ 10-11 B. 2.5 ´ 10-9 C. 4.6 ´ 10-4 D. 0.046 E. 22
28. Using the acid dissociation constants for glycine (Ka1 = 4.6 ´ 10-3 and Ka2 = 2.5 ´ 10-10), calculate the [Z]/[C+] ratio at pH 8.00. (Z = Zwitterion, C+ = fully protonated cation) A. 2.2 ´ 10-6 B. 6.6 ´ 10-4 C. 2.5 ´ 10-2 D. 4.0 ´ 101 E. 4.6 ´ 105
29. Using the acid dissociation constants for glycine (Ka1 = 4.6 ´ 10-3 and Ka2 = 2.5 ´ 10-10), calculate the [Z]/[A-] ratio at pH 7.00. (Z = Zwitterion, A- = deprotonated anion) A. 2.2 ´ 10-5 B. 2.5 ´ 10-3 C. 4.0 ´ 102 D. 4.6 ´ 104 E. 1.8 ´ 107
30. Using the acid dissociation constants for alanine (Ka1 = 5.1 ´ 10-3 and Ka2 = 1.8 ´ 10-10), calculate the [Z]/[A-] ratio at pH 11.50. (Z = Zwitterion, A- = deprotonated anion) A. 6.2 ´ 10-10 B. 5.1 ´ 101 C. 1.8 ´ 10-2 D. 4.6 ´ 109 E. 1.6 ´ 109
31. For serine (Ka1 = 6.5 ´ 10-3 and Ka2 = 6.2 ´ 10-10), at what pH does [Z] = [C+]? A. 2.19 B. 4.79 C. 7.00 D. 9.21 E. 11.81
32. For serine (Ka1 = 6.5 ´ 10-3 and Ka2 = 6.2 ´ 10-10), at what pH does [Z] = [A-]? A. 2.19 B. 4.79 C. 7.00 D. 9.21 E. 11.81
33. For valine (Ka1 = 5.2 ´ 10-3 and Ka2 = 1.9 ´ 10-10), what is the pH at the isoelectric point? A. 2.28 B. 6.00 C. 7.00 D. 8.00 E. 9.72
34. The pH at which the maximum concentration of the zwitterion exists for an amino acid is referred to as the _____ . A. isoelectric point B. zwitterion ion point C. neutral point D. principal point E. deprotonated point
35. When glycine and alanine combine to form a dipeptide, how many structural isomers are possible? A. No structural isomers are possible. B. Two structural isomers are possible. C. Three structural isomers are possible. D. Four structural isomers are possible. E. Six structural isomers are possible.
36. What is the correct structure of alanylserine? Alanine has a CH3 functional group and serine has a CH2OH functional group.
A.
B.
C.
D.
E.
37. What is the correct structure of serylglycine (Ser-Gly)? Serine has a CH2OH functional group and glycine is functionalized with an H atom.
A.
B.
C.
D.
E.
38. Amino acids polymerize in condensation reactions that result in the formation of an amide linkage (or peptide bond) between amino acid molecules. What is a possible dipeptide formed in the reaction of alanine with phenylalanine?
A.
B.
C.
D. E. None of the above
39. How many tripeptides can be made from alanine, serine, and leucine, using each amino acid only once per tripeptide? A. 3 B. 4 C. 6 D. 9 E. 27
40. How many tripeptides can be made from lysine, glycine, and valine, using any number of each amino acid? A. 3 B. 4 C. 6 D. 9 E. 27
41. How many tetrapeptides can be made from serine, alanine, glycine, and lysine, using each amino acid only once per tetrapeptide? A. 4 B. 6 C. 12 D. 24 E. 36
42. How many different dipeptides can be made from the 20 common amino acids? A. 20 B. 39 C. 40 D. 380 E. 400
43. How many different tripeptides can be made from the 20 common amino acids? A. 60 B. 400 C. 8.0 ´ 103 D. 1.6 ´ 105 E. 3.5 ´ 109
44. A certain polypeptide is shown by acid hydrolysis to contain only three amino acids: alanine(Ala), valine(Val), and serine(Ser) with mole fractions of , , and , respectively. Enzymatic hydrolysis yields the following fragments: Ala-Ser, Ala-Val, Val-Ala. What is the primary structure of the polypeptide? A. Ala-Val-Ala-Ser B. Ala-Ser-Ala-Val C. Val-Ala-Ser-Ala D. Ser-Ala-Ala-Val E. Val-Ala-Ala-Ser
45. On complete hydrolysis, a polypeptide gives two glycine, one alanine, one valine, and one serine residue. Partial hydrolysis gives the following fragments: Gly-Ser, Val-Gly, Ser-Val, and Ala-Gly. What is the primary structure of the polypeptide? A. Ala-Gly-Gly-Val-Ser B. Gly-Ser-Val-Gly-Ala C. Ser-Val-Gly-Ala-Gly D. Ala-Gly-Ser-Val-Gly E. Val-Gly-Ser-Val-Ala
46. The enzyme trypsin is used to break bonds in the following polypeptide. Gly-Lys-Met-Leu-Arg-Gly-Met-Arg-Lys What fragments result? A. Gly-Lys-Met; Leu-Arg-Met; Arg-Lys B. Gly-Lys; Met-Leu-Arg; Gly-Met-Arg; Lys C. Gly; Lys-Met-Leu-Arg-Gly; Met-Arg-Lys D. Gly-Lys-Met-Leu-Arg; Gly-Met-Arg-Lys E. Gly-Lys; Met-Leu; Arg-Gly; Met-Arg; Lys
47. The ____ structure corresponds to the amino acid sequence in a molecule. A. nucleic B. primary C. secondary D. tertiary E. quaternary
48. The three-dimensional conformation of a protein is called its A. tertiary structure. B. secondary structure. C. conformational structure. D. primary structure. E. polymerized structure.
49. Watson, Crick and Wilkins received the Nobel prize for A. their work on DNA fingerprinting B. determining the secondary structure of DNA C. determining the identity of the unknown soldier D. determining the primary structure of DNA E. determining the tertiary structure of DNA
50. Diols react with dioic acids to produce A. polyamides and water B. polyesters and water C. condensation polymers and water D. addition polymers E. b and c
Chapter 23--Organic Polymers, Natural and Synthetic Key
1. All of the following statements concerning polyethylene are correct EXCEPT A. plastic bags at supermarkets are made from polyethylene. B. linear polyethylene is more rigid than branched polyethylene. C. linear polyethylene is referred to as high-density polyethylene. D. polyethylene is made by addition reaction of ethylene monomer units. E. the branches in branched polyethylene are primarily carboxylic acid groups.
2. Low density polyethylene (LDPE) is most likely to be used in A. sandwich bags. B. bottle caps. C. rigid pipes. D. contact lenses. E. fabrics and apparel.
3. PVC is a stiff, rugged polymer made from the addition reaction of ____. A. C2H4 B. C3H6 C. CH2CHCl D. CH2CHCN E. C2F4
4. Teflon is made from the addition reaction of ____. A. C2H4 B. C3H6 C. CH2CHCl D. CH2CHCN E. C2F4
5. Polyacrylonitrile is an addition polymer made from the acrylonitrile monomer, CH2CHCN. How many monomer units are in a polyacrylonitrile polymer that has a molar mass of 8.33 ´ 104 g/mol? A. 7.90 ´ 102 B. 1.57 ´ 103 C. 3.14 ´ 103 D. 8.33 ´ 104 E. 4.42 ´ 106
6. Which one of the following polymers is made by condensation reaction? A. Teflon B. poly(ethylene terephthalate) C. polyethylene D. polystyrene E. polyvinyl chloride
7. In principle, the acrylonitrile monomer (CH2CHCN) can polymerize in which of the following ways? 1. 2. 3.
A head-to-tail polymer A head-to-head, tail-to-tail polymer A random polymer
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3 8. The reaction between an amine and a carboxylic acid is a condensation. What small molecule is formed as a byproduct of this reaction? A. H2O B. H2 C. CO2 D. O2 E. CO
9. When a diamine reacts with a dicarboxylic acid, a ____ is formed? A. polyamide B. polyester C. disaccharide D. carbohydrate E. saccharide
10. The polymerization of 1,2-dichloroethylene, CHClCHCl, will yield which of the following polymers?
A.
B.
C.
D.
E.
11. Which of the following molecules will react with oxalic acid to produce a polymer with the following repeating unit?
A.
B.
C.
D.
E.
12. What is the monomer used to make the following polymer?
A.
B.
C.
D.
E.
13. Which of the following statements is/are CORRECT concerning Nylon-66? 1. 2. 3.
The polymer forms at the interface of two insoluble liquds. One monomer is a dicarboxylic acid, the other monomer is a diamine. The polymerization reaction is a condensation with ammonia formed as a byproduct.
A. 1 only B. 2 only C. 3 only D. 1 and 2 E. 1, 2, and 3
14. Glycine, H2NCH2COOH, can polymerize via a condensation reaction. What simple molecule is a byproduct of this polymerization reaction? A. H2 B. NH3 C. N2 D. H2O E. O2
15. The general formula of a carbohydrate is ____. A. CnH2nOn B. CnH2n+2On C. CnH2nO2n D. Cn(H2O)m E. Cn(OH)m
16. Which of the following statements are correct? 1. 2. 3.
Monosaccharides contain either five or six carbon atoms per molecule. Disaccharides are formed when two monosaccharide units combine with the elimination of water. Polysaccharides generally contain anywhere from several hundred to several thousand monosaccharide units.
A. 1 only B. 2 only C. 3 only D. 2 and 3 E. 1, 2, and 3 17. Which of the following statements concerning glucose are correct? 1. 2. 3.
As a pure solid, glucose exists primarily as six-membered ring molecules. In the six-membered ring structure, a CH2OH group is bonded to carbon atom 5. a-glucose and b-glucose are not enantiomers. Although they have several chiral carbon atoms, they differ in configuration at only one carbon atom.
A. 1 only B. 2 only C. 3 only D. 2 and 3 E. 1, 2, and 3 18. Which of the following statements concerning glucose are correct? 1. 2. 3.
a-Glucose and b-glucose each have one chiral carbon atom. In aqueous solution, a-glucose and b-glucose exist as an equilibrium mixture. a-Glucose and b-glucose are enantiomers.
A. 1 only B. 2 only C. 3 only D. 2 and 3 E. 1, 2, and 3 19. Sucrose is commonly known as ____. A. aspartame B. Splenda C. NutraSweet D. saccharine E. sugar
20. Fructose is A. commonly known as blood sugar. B. a monosaccharide found in honey and fruit juices. C. a disaccharide formed from two glucose molecules. D. a polysaccharide formed from several hundred glucose molecules. E. an enzyme that catalyzes the decomposition of glucose.
21. Which of the following is a polysaccharide? A. a-glucose B. fructose C. cellulose D. sucrose E. a-maltose
22. Which of the following statements concerning starch is INCORRECT? A. Starch is a polysaccharide found in plants. B. Starch is abundant in corn and potatoes. C. Amylose comprises about 20% of starch and it is insoluble in water. D. Amylopectin, a highly branched component of starch, is soluble in water. E. Amylose is composed of a-glucose polymers; amylopectin is composed of b-glucose polymers.
23. Which of the following statements concerning cellulose is INCORRECT? A. Humans lack the enzymes needed to catalyze the hydrolysis of cellulose. B. Cellulose contains unbranched chains of roughly 50 - 300 glucose molecules. C. The molecular structure of cellulose allows for hydrogen bonding between polymer chains. D. The major industrial source of cellulose is wood. E. In cellulose, the oxygen bridge between glucose units is in the beta position.
24. Which amino acid contains no chiral carbon atoms? A. proline B. glutamine C. lysine D. glycine E. leucine
25. How many of the 20 common a-amino acids are considered "essential" (i.e., required by humans for protein construction). A. 2 B. 5 C. 10 D. 13 E. 17
26. For an amino acid, proton transfer occurs from an carboxylic acid group to its amine group. The resulting species, which has both a +1 and a -1 charge, is called a(n) _____ . A. amide B. zwitterion C. cation D. anion E. protein
27. Using the acid dissociation constants for glycine (Ka1 = 4.6 ´ 10-3 and Ka2 = 2.5 ´ 10-10), calculate the [Z]/[C+] ratio at pH 1.00. (Z = Zwitterion, C+ = fully protonated cation) A. 2.5 ´ 10-11 B. 2.5 ´ 10-9 C. 4.6 ´ 10-4 D. 0.046 E. 22
28. Using the acid dissociation constants for glycine (Ka1 = 4.6 ´ 10-3 and Ka2 = 2.5 ´ 10-10), calculate the [Z]/[C+] ratio at pH 8.00. (Z = Zwitterion, C+ = fully protonated cation) A. 2.2 ´ 10-6 B. 6.6 ´ 10-4 C. 2.5 ´ 10-2 D. 4.0 ´ 101 E. 4.6 ´ 105
29. Using the acid dissociation constants for glycine (Ka1 = 4.6 ´ 10-3 and Ka2 = 2.5 ´ 10-10), calculate the [Z]/[A-] ratio at pH 7.00. (Z = Zwitterion, A- = deprotonated anion) A. 2.2 ´ 10-5 B. 2.5 ´ 10-3 C. 4.0 ´ 102 D. 4.6 ´ 104 E. 1.8 ´ 107
30. Using the acid dissociation constants for alanine (Ka1 = 5.1 ´ 10-3 and Ka2 = 1.8 ´ 10-10), calculate the [Z]/[A-] ratio at pH 11.50. (Z = Zwitterion, A- = deprotonated anion) A. 6.2 ´ 10-10 B. 5.1 ´ 101 C. 1.8 ´ 10-2 D. 4.6 ´ 109 E. 1.6 ´ 109
31. For serine (Ka1 = 6.5 ´ 10-3 and Ka2 = 6.2 ´ 10-10), at what pH does [Z] = [C+]? A. 2.19 B. 4.79 C. 7.00 D. 9.21 E. 11.81
32. For serine (Ka1 = 6.5 ´ 10-3 and Ka2 = 6.2 ´ 10-10), at what pH does [Z] = [A-]? A. 2.19 B. 4.79 C. 7.00 D. 9.21 E. 11.81
33. For valine (Ka1 = 5.2 ´ 10-3 and Ka2 = 1.9 ´ 10-10), what is the pH at the isoelectric point? A. 2.28 B. 6.00 C. 7.00 D. 8.00 E. 9.72
34. The pH at which the maximum concentration of the zwitterion exists for an amino acid is referred to as the _____ . A. isoelectric point B. zwitterion ion point C. neutral point D. principal point E. deprotonated point
35. When glycine and alanine combine to form a dipeptide, how many structural isomers are possible? A. No structural isomers are possible. B. Two structural isomers are possible. C. Three structural isomers are possible. D. Four structural isomers are possible. E. Six structural isomers are possible.
36. What is the correct structure of alanylserine? Alanine has a CH3 functional group and serine has a CH2OH functional group.
A.
B.
C.
D.
E.
37. What is the correct structure of serylglycine (Ser-Gly)? Serine has a CH2OH functional group and glycine is functionalized with an H atom.
A.
B.
C.
D.
E.
38. Amino acids polymerize in condensation reactions that result in the formation of an amide linkage (or peptide bond) between amino acid molecules. What is a possible dipeptide formed in the reaction of alanine with phenylalanine?
A.
B.
C.
D. E. None of the above
39. How many tripeptides can be made from alanine, serine, and leucine, using each amino acid only once per tripeptide? A. 3 B. 4 C. 6 D. 9 E. 27
40. How many tripeptides can be made from lysine, glycine, and valine, using any number of each amino acid? A. 3 B. 4 C. 6 D. 9 E. 27
41. How many tetrapeptides can be made from serine, alanine, glycine, and lysine, using each amino acid only once per tetrapeptide? A. 4 B. 6 C. 12 D. 24 E. 36
42. How many different dipeptides can be made from the 20 common amino acids? A. 20 B. 39 C. 40 D. 380 E. 400
43. How many different tripeptides can be made from the 20 common amino acids? A. 60 B. 400 C. 8.0 ´ 103 D. 1.6 ´ 105 E. 3.5 ´ 109
44. A certain polypeptide is shown by acid hydrolysis to contain only three amino acids: alanine(Ala), valine(Val), and serine(Ser) with mole fractions of , , and , respectively. Enzymatic hydrolysis yields the following fragments: Ala-Ser, Ala-Val, Val-Ala. What is the primary structure of the polypeptide? A. Ala-Val-Ala-Ser B. Ala-Ser-Ala-Val C. Val-Ala-Ser-Ala D. Ser-Ala-Ala-Val E. Val-Ala-Ala-Ser
45. On complete hydrolysis, a polypeptide gives two glycine, one alanine, one valine, and one serine residue. Partial hydrolysis gives the following fragments: Gly-Ser, Val-Gly, Ser-Val, and Ala-Gly. What is the primary structure of the polypeptide? A. Ala-Gly-Gly-Val-Ser B. Gly-Ser-Val-Gly-Ala C. Ser-Val-Gly-Ala-Gly D. Ala-Gly-Ser-Val-Gly E. Val-Gly-Ser-Val-Ala
46. The enzyme trypsin is used to break bonds in the following polypeptide. Gly-Lys-Met-Leu-Arg-Gly-Met-Arg-Lys What fragments result? A. Gly-Lys-Met; Leu-Arg-Met; Arg-Lys B. Gly-Lys; Met-Leu-Arg; Gly-Met-Arg; Lys C. Gly; Lys-Met-Leu-Arg-Gly; Met-Arg-Lys D. Gly-Lys-Met-Leu-Arg; Gly-Met-Arg-Lys E. Gly-Lys; Met-Leu; Arg-Gly; Met-Arg; Lys
47. The ____ structure corresponds to the amino acid sequence in a molecule. A. nucleic B. primary C. secondary D. tertiary E. quaternary
48. The three-dimensional conformation of a protein is called its A. tertiary structure. B. secondary structure. C. conformational structure. D. primary structure. E. polymerized structure.
49. Watson, Crick and Wilkins received the Nobel prize for A. their work on DNA fingerprinting B. determining the secondary structure of DNA C. determining the identity of the unknown soldier D. determining the primary structure of DNA E. determining the tertiary structure of DNA
50. Diols react with dioic acids to produce A. polyamides and water B. polyesters and water C. condensation polymers and water D. addition polymers E. b and c