Sat chemistry notes by Andora Conti

CHEMISTRY

Dr. D. Bampilis

SAT CHEMISTRY Dr. D. Bampilis

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SAT CHEMISTRY Dr. D. Bampilis

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SAT Chemistry Introduction to Chemistry You have to know: m d V Matter Mixture Homogenous

Heterogeneous

Substance Compound

Element

Chemical - Physical Properties E = m.c2 SI Prefixes T/G/M/k d/c/m/μ/n/p T = θ + 273 Κ = οC + 273 Significant figures The result of a calculation cannot be more accurate than the least accurate number in the calculation.

SAT CHEMISTRY Dr. D. Bampilis

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1. Atomic structure To describe the location of electrons, we use quantum numbers. Quantum numbers are basically used to describe certain aspects of the locations of electrons. For example, the quantum numbers n, l, and ml describe the position of the electron with respect to the nucleus, the shape of the orbital, and its special orientation, while the quantum number ms describes the direction of the electron’s spin within a given orbital. Below are the four quantum numbers, showing how they are depicted and what aspects of electrons they describe. Principal Has positive values of 1, 2, 3, etc. As n increases, the orbital quantum number becomes larger-this means that the electron has a higher (n) energy level and is less tightly bound to the nucleus. Second quantum Has values from 0 to n - 1. This defines the shape of the orbital, number or and the value of l is designated by the letters s, p, d, and f, azimuthal which correspond to values for l of 0, 1, 2, and 3. In other quantum number words, if the value of l is 0, it is expressed as s; if l = 1 = p, l = 2 (l ) = d, and l = 3 = f. Magnetic Determines the orientation of the orbital in space relative to quantum number the other orbitals in the atom. This quantum number has values (ml) from -l through 0 to +l. Spin quantum Specifies the value for the spin and is either +1/2 or -1/2. No number more than two electrons can occupy any one orbital. In order (ms) for two electrons to occupy the same orbital, they must have opposite spins. Orbitals that have the same principal quantum number, n, are part of the same electron shell. For example, orbitals that have n = 2 are said to be in the second shell. When orbitals have the same n and l, they are in the same subshell; so orbitals that have n = 2 and l = 3 are said to be 2f orbitals, in the 2f subshell. Finally, you should keep in mind that according to the Pauli exclusion principle, no two electrons in an atom can have the same set of four quantum numbers. This means no atomic orbital can contain more than two electrons, and if the orbital does contain two electrons, they must be of opposite spin.

You have to know: Proton - Neutron - Electron Atomic number (Z) - Isotopes Mass number (A) Bohr model - atomic spectra Lymar - Balmer - Paschen Mass spectroscopy (vaporizer - ionizer - accelerator - deflector - detector - recorder) The wave - mechanical model • Max Planck • Louis de Broglie • Heisenberg Quantum numbers Principal quantum number (n): K:1, L:2, M:3, N:4 SAT CHEMISTRY Dr. D. Bampilis

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Angular momentum (l) quantum number l = 0……n-1 l: 0 (s),1(p),2(d),3(f) Magnetic quantum number (ml) ml = -l, ….,0, …. +l Spin quantum number (ms) Orbitals Pauli Exclusion Principle Hund’s rule Aufbau Principle Electron configuration

Exercises 1. An element consists of three isotopes in the relative abundance given below. What is the atomic mass of this element? 30.00% = 40.00 amu 50.00% = 41.00 amu 20.00% = 42.00 amu (Α) 40.90 (Β) 41.00 (C) 41.9 (D) 42.20 (Ε) 42.90 2. The total number of electrons that can be accommodated in the fourth principal energy level is: (Α) 2 (Β) 8 (C) 18 (D) 32 (Ε) 50 3. If the set of quantum numbers n = 3, l = 1, ml = 0, ms = +1/2 represents the last electron to be added to complete the ground state electron configuration of an element, which one of the following could be the symbol for the element? (A) Na (B) Si (C) Th (D) V (E) Zn 4. Which element has the electron configuration:1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 ? (Α) Cr (Β) Μn (C) Μo (D) S (Ε) Se 5. Which of the following elements has electrons in f orbitals? (Α) Ar (Β) O (C) S (D) Ti (Ε) U 6. Which of the following elements has the electron configuration: 1s 2 2s 2 2p 6 3s 2 3p 4 ? (Α) Ar (Β) O (C) S (D) Ti (Ε) U 7. Which of the following elements has the same number of electrons as Ca2+? (Α) Ar (Β) O (C) S (D) Ti (Ε) U 8. Which gas-phase atom in its ground state could have an electron with the quantum numbers: n = 3, l = 2, ml = 0, ms = -1/2? (A) Na (B) Mg (C) P (D) Ti 9. Which set of quantum numbers could represent an electron in a 5f orbital? SAT CHEMISTRY Dr. D. Bampilis

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(A) l = 4, ml = 2

(B) l = 2, ml = -3

(D) l = 3, ml = 0

10. Which set of quantum numbers (n, l, ml, ms) is possible for the outermost electron in a strontium atom in its ground state? (A) 5, 0, 0, -1/2 (B) 5, 0, 1, 1/2 (C) 5, 1, 0, 1/2 (D) 5, 1, 1, -1/2 11. Which quantum numbers represent the orbitals being filled in the ground state for the elements Sc (21) to Zn (30)? (A) n = 3, l = 1 (B) n = 3, l = 2 (C) n = 4, l = 1 (D) n = 4, l = 2 12. Which set of quantum numbers (n, l, ml) is forbidden? (A) 3, 2, 0 (B) 3, 1, –1 (C) 2, 0, 0

(D) 1, 1, 0

13. Which characteristic of an atomic orbital is most closely associated with the magnetic quantum number, ml? (A) size (B) shape (C) occupancy (D) orientation

SAT CHEMISTRY Dr. D. Bampilis

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2. Periodicity You have to know: Electron configuration and Periodic Table s, p, d, f block Transition elements - d block d: half - filled or full - filled Periodic law Alkali metals (1) - Alkaline Earth Metals (2) - Halogens (17) - Noble gases (18) - Metals - Nom metals - Metalloids (B, Si, Ge, As, Sb, Te) Lanthanides - Actinides: f block Atomic Radii Ionic radius Electronegativity Electron affinity Ionization Energy

Why is fluorine out of line? The incoming electron is going to be closer to the nucleus in fluorine than in any other of these elements, so you would expect a high value of electron affinity. However, because fluorine is such a small atom, you are putting the new electron into a region of space already crowded with electrons and there is a significant amount of repulsion. This repulsion lessens the attraction the incoming electron feels and so lessens the electron affinity. A similar reversal of the expected trend happens between oxygen and sulfur in Group 6. The first electron affinity of oxygen (-142 kJ.mol-1) is smaller than that of sulfur (-200 kJ.mol-1) for exactly the same reason that fluorine's is smaller than chlorine's.

Second electron affinity You are only ever likely to meet this with respect to the group 6 elements oxygen and sulfur which both form -2 ions. The second electron affinity is the energy required to add an electron to each ion in 1 mole of gaseous 1- ions to produce 1 mole of gaseous 2- ions. This is more easily seen in symbol terms. X−(g) + e−

X2−(g)

It is the energy needed to carry out this change per mole of X−. SAT CHEMISTRY Dr. D. Bampilis

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Why is energy needed to do this? You are forcing an electron into an already negative ion. It's not going to go in willingly! O(g) + e-

O-(g)

1st EA = -142 kJ.mol-1

O-(g) + e-

O2-(g)

2nd EA = +844 kJ.mol-1

The positive sign shows that you have to put in energy to perform this change. The second electron affinity of oxygen is particularly high because the electron is being forced into a small, very electron-dense space.

Exercises 1. Which of the following atoms would have the largest second ionization energy? (Α) Mg (Β) Cl (C) S (D) Ca (Ε) Νa 2. Order the elements S, CI, and F in terms of increasing atomic radii. (Α) S, CI, F (Β) CI, F, S (C) F, S, CI (D) F, CI, S (Ε) S, F, CI 3. Which of the following elements is the LEAST chemically reactive? (Α) Ar (Β) O (C) S (D) Ti (Ε) U 4. For elements in the left-most column of the periodic table, properties that have increasing values as the atomic number increases include which of the following? I. Ionization energy (potential) II. Atomic radius III. Atomic mass (A) I only (B) III only (C) I and II only (D) II and III only (E) I, II, and III 5. Removing an electron from so dium is an ..... process and removing an electron from fluorine is an ..... process. (Α) Endothermic, exothermic (Β) Exothermic, endothermic (C) Endothermic, endothermic (D) Exothermic, exothermic (Ε) More information is needed For questions 6 - 10: Statement I BECAUSE Statement II 6. The second ionization energy of Β is higher than that of Be. BECAUSE The second electron to be removed from Β and Be comes from the same principal energy level. 7. Oxygen has a smaller first ionization energy than fluorine. BECAUSE Oxygen has a higher Ζ eff value than does fluorine. 8. Hydrogen has a lower ionization energy than does helium. BECAUSE Hydrogen bonds with halogens to form polar covalent bonds. SAT CHEMISTRY Dr. D. Bampilis

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9. Potassium has a lower first ionization energy than lithium has. BECAUSE Potassium has more protons in its nucleus than lithium has. 10. An element that has the electron configuration 1s 2 2s 2 2p 6 3s 2 3p 6 3d 3 4s 2 is a transition element. BECAUSE In atoms of transition elements, the 1s, 2s, 2p, 3s and 3p orbitals are completely filled in the ground state.

Atomic structure & periodicity For questions 1 - 5: (Α) Alkali metals (D) Halogens

(Β) Alkaline Earth metals (E) Rare earth metals

1. Which of the following is used primarily in semiconductors? 2. Which occur as diatomics? 3. Which make oxides with the formula X2O? 4. Which have large electronegativity values? 5. Which have small ionization energies? For questions 6 - 9: (Α) Na+ (B) Al

(D) Ti

(E) B

6. Which has seven valence electrons? 7. Which has an electron configuration 1s22s22p63s23p1? 8. Which has the same electron configuration as the neon atom? 9. Which has valence electrons in the d orbitals? For questions 10 - 14: (Α) Alkali metals (D) Halogens

(B) Alkaline Earth metals (E) Transition metals

10. Which is the most unreactive family of elements? 11. Which form negative ions in an ionic bond? SAT CHEMISTRY Dr. D. Bampilis

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12. Which consist of atoms that have valence electrons in a d subshell? 13. Which exist as diatomic molecules at room temperature? 14. Which group possesses the lowest first ionization energy in their respective period? For questions 15 - 17: (Α) Bohr model (B) De Broglie’s wave hypothesis (C) Heisenberg’s uncertainty principle (D) Quantum theory

(E) Atomic theory

15. Which principle provides that all matter may be considered a wave? 16. What views electrons in true orbits around the nucleus? 17. What considers that one cannot know position and velocity of an electron at the same moment? For questions 18 - 32: Statement I BECAUSE Statement II 18. The metalloids have similar characteristics BECAUSE Their valence shells have the same configuration 19. Metals are good conductors of heat and electricity BECAUSE The positive nuclei are surrounded by a “sea” of mobile electrons 20. Elements in a group have similar properties BECAUSE Their valence shells have the same energy 21. The first ionization energy for an atom is greater than the second ionization energy BECAUSE The closer an electron is to the nucleus, the more difficult it is to remove 22. Sodium has a smaller atomic radius than chlorine BECAUSE A sodium atom does not have as many valence electrons as a chlorine atom does. 23. Carbon’s electric configuration is 1s22s22p2 rather than 1s22s23s2 BECAUSE 3s electrons are lower in energy than 2p electrons 24. The properties of phosphorus should be closer to those of sulfur than to those of nitrogen SAT CHEMISTRY Dr. D. Bampilis

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BECAUSE Phosphorus and nitrogen are in the same row of the periodic table 25. The halogens, in group VIIA, all form stable diatomic molecules BECAUSE They each need one electron to fill their outer shells 26. Metals are good conductors of electricity BECAUSE They are held together by ionic bonds 27. Two electrons in the 2s subshell must have opposite spins BECAUSE The Pauli exclusion principle states that no two electrons in the same atom can have identical quantum numbers 28. 40Ca is a neutral atom BECAUSE It has the same number of protons and neutrons 29. The most important factor in determining the chemical properties of an element is the number of electrons in the outermost shell BECAUSE The number of electrons in the outermost shell determines the bonding characteristics of the element 30. Iron is an element BECAUSE It cannot be broken into smaller units and retains its physical and chemical properties 31. An element (X) with an atomic number of 16 has 14 electrons in X2+ BECAUSE It has gained two electrons 32. Atomic radii increase down a group BECAUSE The higher the atomic number within a group, the smaller the atom 33. The element with atomic number 32 describes: (Α) A metal (B) A non-metal (C) A metalloid

(D) A halogen (E) A noble gas

34. How many neutrons are probably in the nucleus of an element of atomic weight 197? (Α) 43 (B) 79 (C) 83 (D) 100 (E) 118 35. The transition metals are characterized by: SAT CHEMISTRY Dr. D. Bampilis

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(A) completely filled d subshells (C) partially filled d subshells (E) both (a) and (c) are correct

(B) completely filled f subshells (D) partially filled f subshells

36. Neutral atoms of F (fluorine) have the same number of electrons as: (A) B3(B) N+ (C) Ne(D) Na(E) Mg3+ 37. Which of the elements in Group 1A of the periodic table has the greatest metallic character? (A) H (B) Li (C) Na (D) K (E) Rb 38. The ionization energy of an element is: (A) a measure of its mass (B) the energy required to remove an electron from the element in its gaseous state (C) the energy released by the element in forming an ionic bond (D) the energy released by the element upon receiving an additional electron (E) none of the above 39. Elements in a row have the same: (A) Atomic weight (C) Maximum principal quantum number (n) (E) Atomic number

(B) Maximum azimuthal quantum number(l) (D) Valence electron structure

40. Which of the following has the largest radius? (A) Sr (B) P (C) Mg (D) Al3+

(E) Mg2+

41. Which of the following elements has the lowest electronegativity? (A) Cesium (B) Strontium (C) Calcium (D) Barium (E) Potassium 42. Which of the following is biggest in size? (A) Ca (B) Ca+ (C) Ca2+

(D) Ca-

(E) Ca2-

43. The order of the elements in the periodic table is based on: (A) the number of neutrons (B) the radius of the atom (C) the atomic number (D) the atomic weight (E) the number of oxidation states 44. The elements within each column of the Periodic Table: (A) have similar valence electron configurations (B) have similar atomic radii (C) have the same principal quantum number (D) will react to form stable elements (E) have no similar chemical properties 45. Which of the following has the highest 1st ionization energy? (A) Ga (B) Ba (C) Ru (D) F (E) N 46. Which of the following has the lowest electronegativity? SAT CHEMISTRY Dr. D. Bampilis

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(A) Ca

(B) Cl

(D) P

(E) Zn

47. Which element has the greatest electronegativity? (A) Chlorine (B) Oxygen (C) Sulfur (D) Phosphorus (E) Fluorine 48. Transition metal compounds generally exhibit bright colors because: (A) The electrons in the partially filled d orbitals are easily promoted to excited states (B) The metals become complexed with water (C) The metals conduct electricity, producing colored light (D) The electrons in the d orbitals emit energy as they relax (E) Their valence electrons cause them to bind to other metals 49. Which of the following is a non metal? (A) Fr (B) Pd (C) I

(D) Sc

(E) Sr

50. Which of the following has the greatest affinity for electrons? (A) F (B) Cl (C) Br (D) K (E) C 51. Which of the following is the most electronegative element? (A) He (B) I (C) N (D) O (E) C 52. Which of the following is not a property of Group IA elements? (A) Low ionization energies (B) Low electronegativities (C) High melting points (D) Metallic bonding (E) Electrical conductivity 53. Arrange the following elements in order of decreasing nonmetallic character: Ge, Sn, Pb, Si. (A) Pb, Sn, Ge, Si (B) Ge, Sn, Pb, Si (C) Si, Ge, Sn, Pb (D) They all have equal nonmetallic character since they are all in the same column of the Periodic Table (E) None of the above 54. Electron affinity is defined as: (A) the change in energy when a gaseous atom in its ground state gains an electron (B) the pull an atom has on the electrons in a chemical bond (C) the energy required to remove a valence electron from a neutral gaseous atom in its ground state (D) the energy difference between an electron in its ground state and its excited state (E) none of the above 55. Which of the following is an incorrect association? (A) Mendeleev-periodic table (B) Faraday-electrolytic cells (C) Millikan-charge of electrons (D) Rutherford-photoelectric effect (E) They are all correct

SAT CHEMISTRY Dr. D. Bampilis

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56. Members of group 1 have similar reactivity because they have: (A) the same number of protons (B) the same number of electrons (C) similar outer shell configurations (D) valence electrons with the same quantum numbers (E) the same number of neutron 57. Boron found in nature has an atomic weight of 10.811 and is made up of the isotopes 10B (mass 10.013 amu) and 11B (mass 11.0093 amu). What percentage of naturally occurring boron is made up of 10B and 11B respectively? (A) 30:70 (B) 25:75 (C) 20:80 (D) 15:85 58. The modern periodic table is ordered on the basis of: (A) atomic mass (B) atomic radius (C) atomic charge (D) atomic number (E) number of neutron 59. The electron configuration 1s22s22p63s23p64s23d7 represents an atom of the element: (A) Br (B) Co (C) Cd (D) Ga (E) Mg 60. A neutral atom whose electron configuration is 1s22s22p63s23p64s23d104p65s24d105p6 is: (A) Highly reactive (B) A noble gas (C) A positively charged ion (D) A transition metal (E) A lanthanide element

SAT CHEMISTRY Dr. D. Bampilis

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3. Bonding You have to know: Ionic ΔΕ ≥ 1.7 Covalent (polar - non polar) - Lewis structure Exceptions to the octet rules: BeH2, BF3, BH3, PCl5, SCl6 Coordinate Covalent bonds Metallic Intermolecular forces (Van der Walls - SAT) • Dipole - Dipole attraction • London Dispersion Forces (Van der Walls - IB) • Hydrogen Bonds Molecule - ion attraction Resonance structure VSEPR polar - non polar molecules Hybridization Sigma - pi bonds

Ionic substances Solid: don’t conduct electric current liquid: conduct electric current Melting Point Boiling Point: high Low volatilities and vapor pressures Brittle solids Soluble in water (with exceptions)

Properties Molecular crystals and liquids Solid or liquid: don’t conduct electric current Melting Point – Boiling Point: Low Relative volatile Soft - waxy (solids) Large amount of energy to decompose the substance

Exercises For questions 1 - 4: (A) an ionic substance (B) a polar covalent substance (C) a nonpolar covalent substance (D) an amorphous substance (E) a metallic network 1. KCl(s) is: 2. HCl(g) is: 3. CH4(g) is: 4. Li(s) is: For questions 5 - 8: (A) hydrogen bond (E) metallic bond

(B) ionic bond (C) polar covalent bond (D) pure covalent bond

SAT CHEMISTRY Dr. D. Bampilis

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5. The type of bond between atoms of potassium and chloride in a crystal of potassium chloride is: 6. The type of bond between the atoms in a nitrogen molecule is: 7. The type of bond between atoms in a molecule of CO2 (electronegativity difference = ~1) is: 8. The type of bond between atoms of calcium in a crystal of calcium is: For questions 9 - 11: (A) zero

(B) one

(D) three

(E) four

9. The number of bonds predicted for O2. 10. The number of bonds predicted for N2. 11. The number of bonds predicted for H2. For questions 12 - 15: (A) Linear geometry (D) Pyramidal geometry 12. NH3 has a:

(B) Bent geometry (C) Tetrahedral geometry (E) Equilateral triangle geometry

13. H2O has a: 14. BeF2 has a: 15. CH4 has a: For questions 16 - 18: (A) BeF2

(B) NH3

(D) CH2CH2

(E) CCl4

16. This species has sp2 hybrid orbitals. 17. This species has sp hybrid orbitals. 18. This species contains a pi bond. For questions 19 - 22: (A) hydrogen bonding (B) ionic bonding (C) metallic bonding (D) nonpolar covalent bonding (E) polar covalent bonding 19. This holds a sample of barium iodide, BaI2, together.

SAT CHEMISTRY Dr. D. Bampilis

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20. This allows many solids to conduct electricity. 21. This attracts atoms of hydrogen to each other in a H2 molecule. 22. This is responsible for the relatively high boiling point of water. For questions 23 - 32: Statement I BECAUSE Statement II 23. Nonmetallic atoms of the same element combine covalently BECAUSE The two elements have the same electronegativities 24. A nonpolar molecule can have polar bonds BECAUSE Polar bonds can be symmetrically arranged in a molecule so that there are no net poles 25. The bond in an O2 molecule is considered to be nonpolar BECAUSE The oxygen atoms in an O2 molecule share the bonding electrons equally 26. An ionic solid is a good conductor of electricity BECAUSE An ionic solid is composed of positive and negative ions joined together by electrostatic forces 27. The hybrid orbitals of carbon in acetylene are believed to be the sp form BECAUSE Acetylene is a linear compound with a triple bond between the carbons 28. Atom A with 7 valence electrons forms AB2 with atom B with two valence electrons BECAUSE B donates its electrons to fill the outer shell of A 29. Water is a polar substance BECAUSE The bonding electrons in water are shared equally 30. He2 is not known to commonly form BECAUSE He is lighter than air 31. CCl4 is a nonpolar molecule BECAUSE The dipole moments in CCl4 cancel each other out

SAT CHEMISTRY Dr. D. Bampilis

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32. One of the most important factors in determining the chemical properties of an element is the number of electrons in its outermost shell BECAUSE The number of electrons in the outer shell determines the bonding characteristics of an element 33. An sp2 configuration is represented by which orientation? (A) Tetrahedral (B) Planar (C) Linear (D) Trigonal planar

(E) Square

34. When the electrons are shared unequally by two atoms, the bond is said to be: (A) covalent (B) polar covalent (C) coordinate covalent (D) ionic (E) metallic 35. Which of the following contains a coordinate covalent bond? (A) HCl (Β) H2O (C) H2 (D) H3O+

(E) NaCl

36. Which of the following elements can form bonds with sp3 hybridization? (A) Sodium (B) Nitrogen (C) Carbon (D) Oxygen (E) Fluorine 37. A triple bond may be best described as: (A) two sigma bonds and one pi bond (C) two sigma bonds and two pi bonds (E) three pi bonds

(B) one sigma bond and two pi bonds (D) three sigma bonds

38. Molecules of sodium chloride: (A) display ionic bonding (B) display polar covalent bonding (C) are polar (D) dissociate in water solution

(E) do not exist

39. Which of the following molecules is polar? (A) BH3 (B) NF3 (C) C2H6 (D) SF6

(E) CCl4

40. Which of the following molecules has a trigonal pyramidal geometry? (A) BH3 (B) H2O (C) CH4 (D) NH3 (E) AlCl3 41. The shape of a PCl3 molecule is described as: (A) bent (B) trigonal pyramidal (C) linear (D) trigonal planar

(E) tetrahedral

42. The structure of BeCl2 can best be described as: (A) linear (B) bent (C) trigonal (D) tetrahedral

(E) square

43. All of the following have covalent bonds EXCEPT: (A) HCl (B) CCl4 (C) H2O (D) CsF

(E) CO2

44. The complete loss of an electron of one atom to another atom with the consequent formation of electrostatic charges is said to be: (A) A covalent bond (B) A polar covalent bond (C) An ionic bond SAT CHEMISTRY Dr. D. Bampilis

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(D) A coordinate covalent bond

(E) A pi bond between p orbitals

45. Which molecule is incorrectly matched with the molecular geometry? Molecule Molecular geometry (A)

SF6

Octahedral

(B)

CH4

Tetrahedral

(C)

SO 3

trigonal planar

(D)

SeCl 4

tetrahedral

(E)

PH 3

trigonal pyramidal

46. From their electron configurations, one can predict that the geometric configuration for which of the following molecules is NOT correct? (A) PF3 trigonal planar (B) CF4 tetrahedral (C) CHCl3 irregular tetrahedron (D) OF2 bent (v-shaped) (E) HF linear 47. Which numbered response lists all the molecules below that exhibit resonance and none that do not? Ι. AsF5 ΙΙ. ΗΝO 3 III. SO 2 (A) I only (B) II only (C) ΙΙ and III (D) III and IV (Ε) Ι, ΙΙ, and III 48. Which one of the following molecules is octahedral? (Α) BeCl 2 (Β) SeF6 (C) BF3 (D) PF5 (Ε) CF4 49. The sulfur hexafluoride molecule is nonpolar and contains no lone (unshared) electron pairs on the sulfur atom. Which answer choice lists all of the bond angles contained in sulfur hexafluoride? (Α) 120 o (B) 180 o (C) 90 o and 180 o (D)90 o , 120 o , and 180 o (Ε)109.5 o For questions 50 - 54: Statement I BECAUSE Statement II 50. Molecules that contain a polar bond are not necessarily polar compounds. BECAUSE If polar bonds in a molecule are symmetrically arrange d, then their polarities will cancel and they will be nonpolar. 51. The ΝΗ 3 molecule is more polar than the NF 3 molecule. BECAUSE Fluorine atoms are larger than hydrogen atoms. 52. Ice, unlike most substances, is denser than water in the liquid phase. BECAUSE In water, hydrogen bonds can form between the positively charged Η atom on one water molecule and the slightly negatively charged Ο atom on a nearby water molecule. SAT CHEMISTRY Dr. D. Bampilis

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53. Diamond has a high melting point. BECAUSE In a diamond crystal, the carbon atoms are held in place by ionic bonds. 54. A molecule of silicon tetrachloride, SiCl4 , is nonpolar. BECAUSE The four bonds in SiCl4 are identical and the molecule has a tetrahedral structure.

SAT CHEMISTRY Dr. D. Bampilis

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4. Gases and Gas Laws Graham’s Law of Diffusion and Effusion Graham’s law states that the rates of effusion of two gases are inversely proportional to the square roots of their molar masses at the same temperature and pressure:

Effusion is the term used to describe the passage of a gas through a tiny orifice into an evacuated chamber, as shown in the figure below.

The rate of effusion measures the speed at which the gas travels through the tiny hole into a vacuum. Another term to remember for the test is diffusion. Diffusion is the term used to describe the spread of a gas throughout a space or throughout a second substance.

You have to know: Kinetic Molecular Theory Maxwell - Boltzmann distribution Graham’s law of diffusion (effusion)

Rate A Rate B Charles’s law

MB MA V T

Boyle’s law P · V = k

P k T P V P2 V2 Combined gas law 1 1 T1 T2 Gay - Lussac’s law

Dalton’s law of Partical Pressures Ptot = P1 + P2 + P3

n1 Ptot n tot

Ideal gas law: P · V = n·RT (real gases at low pressure and high temperature) Two devices to measure pressure: • Mercury Barometer (Eudiometer) • Manometer SAT CHEMISTRY Dr. D. Bampilis

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Exercises 1. Which of the following gases would be the densest at standard temperature and pressure? (Α) Helium (Β) Argon (C) Carbon dioxide (D) Xenon (Ε) Nitrogen 2. A gas diffuses one-third as fast as O2 at 100oC. This gas could be: (A) He (M = 4) (B) C2H5F (M = 48) (C) C7H12 (M = 96)

(D) C5F12 (M = 288)

3. Ar and He are both gases at room temperature. How do the average molecular velocities (V) of their atoms compare at this temperature? (A) VHe = 10VAr (B) VAr = 10VHe (C) VHe = 3VAr (D) VAr = 3VHe 4. Moist air is less dense than dry air at the same temperature and barometric pressure. Which is the best explanation for this observation? (A) H2O is a polar molecule but N2 and O2 are not. (B) H2O has a higher boiling point than N2 or O2. (C) H2O has a lower molar mass than N2 or O2. (D) H2O has a higher heat capacity than N2 or O2. 5. Two samples of gas, one of argon and one of helium, have the same pressure, temperature and volume. Which statement is true assuming both gases behave ideally? (A) The helium sample contains more atoms than the argon sample and the helium atoms have a higher average speed. (B) The two samples have the same number of atoms but the helium atoms have a higher average speed. (C) The two samples have the same number of atoms and both types of atoms have the same average speed. (D) The two samples have the same number of atoms but the argon atoms have a higher average speed. 6. Which of the following best illustrates a graph of pressure versus volume for a gas at constant temperature?

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7. The bulb of the open-end manometer shown below contains a gas. True statements about this system include which of the following?

I. Only atmospheric pressure is exerted on the exposed mercury surface in the right side of the tube. II. The gas pressure is greater than atmospheric pressure. III. The difference in the height, h, of mercury levels is equal to the pressure of the gas. (A) II only (B) III only (C) I and II only (D) I and III only (E) I, II, III For question 8: Statement I BECAUSE Statement II 8. Statement I: At the same temperature and pressure, 1 L of hydrogen gas and 1 L of neon gas have the same mass. BECAUSE Statement II: Equal volumes of ideal gases at the same temperature and pressure contain the same number of moles.

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5. Stoichiometry You have to know: Mole - Avogadro’s number Avogadro’s law

n V

Molar mass

m Mr

m V

M Vm

P V

m RT M

V Vm

V (STP) n 22, 4

N NA

M (STP) 22, 4

P M d RT

Balancing chemical equations Mole ratio Limiting and excess reagents Percent yield of a product

Exercises For questions 1 - 3: (A) N2O5 (B) N2O3

(D) NO

(E) N2O

1. What is the empirical formula for a compound containing 63.8% N and 36.2% O? 2. What is the empirical formula for a compound containing 36.7% N and 63.3% O? 3. What is the empirical formula for a compound containing 25.9% N and 74.1% O? For questions 4 - 6: (A) 2.294 (B) 36.51

(D) 25.3

(E) 2.513

4. For 4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g), if you begin with 16.00 g ammonia and excess oxygen, how many grams of water will be obtained? 5. For 4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g), if you begin with 66.00 g ammonia and 54.00 g oxygen, how many grams of water will be obtained? 6. For 4NH3(g) + 5O2(g) 2.513 moles of NO?

4NO(g) + 6H2O(g), how many moles of NH3 are needed to produce

For questions 7 - 10: (A) 1.807 x 10-24 (B) 3.476 x 10-2 (C) 1.171 x 10-2 (D) 1.204 x 1024

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(E) 2.414 x 10-1

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7. How many phosphine molecules are in two moles of phosphine? 8. How many moles of CO2 are in 1.53 g CO2? 9. How many atoms are in one mole of water? 10. How many moles are in 4.35 grams of water? 11. When the following is balanced, C4H10 + O2 CO2 + H2O, what is the coefficient of CO2? (A) 2 (B) 4 (C) 8 (D) 10 (E) 13 12. What is the approximate percentage composition by mass of the element oxygen in the compound HClO4? (A) 16% (B) 35% (C) 50% (D) 64% (E) 75% 13. When the following equation is balanced, how many moles of NF3 would be required to react completely with 6 moles of H2O? ___NF3(g) + ___H2O(g) ___HF(g) + ___NO(g) + ___NO2(g) (A) 0.5 mole (B) 1 mole (C) 2 moles (D) 3 moles (E) 4 moles 14. For the following equation, Fe2O3(s) + 3CO(g) 2Fe(s) + 3CO2(g), when 3.0 mol Fe2O3 is allowed to completely react with 56 g CO, approximately how many moles of iron, Fe, are produced? (A) 0.7 (B) 1.3 (C) 2.0 (D) 2.7 (E) 6.0 15. What is the percent by mass of silicon in a sample of SiO2? (A) 21% (B) 33% (C) 47% (D) 54% 16. When the following equation is balanced, ___PH3 + ___O2 the coefficient of H2O? (A) 1 (B) 2 (C) 3 (D) 4

(E) 78% ___P2O5 + ___H2O, what is (E) 5

17. What are the products of the following reaction: H2SO4(aq) + Ba(OH)2(aq) ? (A) O2 (B) BaSO4 (C) O2 and BaSO4 (D) O2 and BaSO4 (E) H2O and BaSO4 18. For the equation, 2Mg(s) + O2(g) 2MgO(s), if 48.6 g Mg is placed in a container with 64.0 g O2 and the reaction is allowed to go to completion, what is the mass of MgO(s) produced? (A) 15.4 g (B) 32.0 g (C) 80.6 g (D) 96.3 g (E) 112 g 19. For the equation, 2NO(g) + 2H2(g) N2(g) + 2H2O(g), which of the following is true? (A) If 1 mole of H2 is consumed, 0.5 moles of N2 is produced (B) If 1 mole of H2 is consumed, 0.5 moles of H2O is produced (C) If 0.5 moles of H2 is consumed, 1 mole of N2 is produced (D) If 0.5 moles of H2 is consumed, 1 mole of NO is produced SAT CHEMISTRY Dr. D. Bampilis

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(E) If 0.5 moles of H2 is consumed, 1 mole of H2O is produced 20. Which of the following expressions is equal to the number of iron (Fe) atoms present in 10.0 g Fe? (atomic mass of Fe = 55.9) (A) 10 x 55.9 x (6.022 x 1023) atoms (B) (6.022 x 1023) / 10 x 55.9 atoms (C) 10 x (6.022 x 1023) / 55.9 atoms (D) 55.9 / 10 x (6.022 x 1023) atoms (E) 10 / (55.9 x 6.022 x 1023) atoms 21. The formula Cr(NH3)5SO4Br represents: (A) 4 atoms (B) 8 atoms (C) 12 atoms

(D) 23 atoms

(E) 27 atoms

22. What is the molecular formula of a compound made of 25.9% N and 74.1% O? (A) NO (B) NO2 (C) N2O (D) N2O5 (E) N2O4 23. The balanced molar relationship from the reaction H2O2 (A) 1:1:1 (B) 2:1:1 (C) 1:2:1 (D) 2:2:1

H2O + O2 is (E) 2:1:2

24. What volume of H2O is required to produce 5 L O2 by the following equation: H2O(g) H2(g) + O2(g)? (A) 3 L (B) 5 L (C) 10 L (D) 16 L (E) 14 L 25. What is the molecular weight of HClO4? (A) 52.5 (B) 73.5 (C) 96.5

(D) 100.5

26. Which of the following molecules contains 17 atoms? (A) Al2(SO4)3 (B) Al(NO3)3 (C) Ca(HCO2)2 (D) Mg(IO3)2

(E) 116.5

(E) Two of the above

27. Twenty liters of NO gas react with excess oxygen. How many liters of NO2 gas are produced if the NO gas reacts completely? (2NO + O2 2NO2) (A) 5 L (B) 10 L (C) 20 L (D) 40 L (E) 50 L 28. How much reactant remains if 92 g HNO3 reacts with 24 g LiOH assuming a complete reaction? (A) 46 g HNO3 (B) 29 g HNO3 (C) 12 g HNO3 (D) 2 g LiOH (E) 12 g LiOH 29. What is the density, at STP, of a diatomic gas whose gram-formula mass is 80 g/mol? (A) 1.9 g/L (B) 2.8 g/L (C) 3.6 g/L (D) 4.3 g/L (E) 5.0 g/L 30. How many liters of H2 can be produced at STP by the decomposition of 3 mol NH3? (A) 4.5 L (B) 27 L (C) 67.2 L (D) 96 L (E) 101 L 31. How many mol CO2 molecules are represented by 1.8 x 1024 atoms? (A) 1 (B) 2 (C) 3 (D) 4 (E) 5 32. How many grams of Na2SO4 can be produced by reacting 98 g H2SO4 with 40 g NaOH? SAT CHEMISTRY Dr. D. Bampilis

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(A) 18 g

(B) 36 g

(D)142 g

33. What are the missing products of the following reaction? NH4Cl + Ca(OH)2 _____ + CaCl2 (A) N2 (B) NH3 (C) H2O (D) NH3 + N2

(E) 150 g

(E) NH3 + H2O

34. How many grams of water can be produced when 8 g of hydrogen react with 8 g oxygen? (A) 8 g (B) 9 g (C) 18 g (D) 27 g (E) 30 g 35. How many atoms are represented in Na2CO3•10H2O (A) 4 (B) 16 (C) 36 (D) 60

(E) 96

36. What is the density of bromine vapor at STP? (A) 2.5 g/L (B) 2.9 g/L (C) 3.6 g/L (D) 4.9 g/L

(E) 7.1 g/L

37. Fill in the missing reactant: NaOH + _____ (A) Cl2 (B) HCl (C) HClO

NaClO2 + H2O (D) HClO2 (E) HClO3

38. How many grams of Na are present in 30 g NaOH? (A) 10 g (B) 15 g (C) 17 g (D) 20 g

(E) 22 g

39. What is the sum of the coefficients when the following reaction is balanced? ___C6H6 + ___O2 ___CO2 + ___H2O? (A) 7 (B) 14 (C) 28 (D) 35 (E) 42 40. How many atoms are represented by the following formula? K3Fe(CN)6 (A) 6 (Β) 10 (C) 16 (D) 20 (E) 18 41. Twenty-two grams of CO2 at STP is identical to (A) 1 mole of CO2 (B) 6.022 x 1023 atoms (D) 11.2 liters (E) 22.4 liters

42. What volume does 8.5 g NH3 occupy at STP? (A) 2.81 L (B) 5.61 L (C) 11.21 L (D) 22.41 L

(E) 44.81 L

43. What is the formula of a hydrocarbon composed of 86% carbon and 14% hydrogen by weight? (A) CH4 (B) C2H4 (C) C2H6 (D) C3H8 (E) C4H6 44. How many grams of CO2 are produced by the complete reaction of 100 g CaCO3 with excess HCl? (A) 22 g (B) 44 g (C) 79 g (D) 110 g (E) 132 g 45. 28 mL of nitrogen are reacted with 15 mL of hydrogen. How many milliliters of which gas are left unreacted? SAT CHEMISTRY Dr. D. Bampilis

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(A) 5 mL H2

(B) 5 mL N2

(D) 11 mL N2

(E) 23 mL N2

46. If 28 mL of nitrogen are reacted with 15 mL of hydrogen, what is the total volume of gas present after the reaction has occurred, assuming volumes are additive? (A) 11 mL (B) 17 mL (C) 27 mL (D) 33 mL (E) 42 mL 47. What is the mass of 1 L of a gas at STP whose molar mass is 254 g/mol? (A) 11.3 g (B) 25.4 g (C) 30.6 g (D) 76.5 g (E) 254 g 48. A type of ion found in aluminum oxide is: (Α) X + (Β) X 2+ (C) X 3+

(D) XO 3 2-

(Ε) XO 4 2-

49. A type of ion found in potassium phosphate is: (Α) X + (Β) X 2+ (C) X 3+ (D) XO 3 2-

(Ε) XO 4 2-

50. A type of ion found in sodium acetate is: (Α) X + (Β) X 2+ (C) X 3+

(Ε) XO 4 2-

(D) XO 3 2-

51. How many moles of hydrogen sulfide are contained in a 35.0 g sample of this gas? (Α) 1.03 mol (Β) 2.06 mol (C) 6.18 mol (D) 9.45 mol (Ε) 11.3 mol 52. What is the molar mass of ethanol (C 2 H5 OH)? (Α) 34.2 (Β) 38.9 (C) 46.1 (D) 45.1

(Ε) 62.1

53. Ammonia can be produced by the reaction of nitrogen and hydrogen gas. Suppose the reaction is carried out starting with 14 g of nitrogen and 15 g of hydrogen. How many grams of ammonia can be produced? (Α) 17.04 g (Β) 34.08 g (C) 51.1 g (D) 85.2 g (Ε) 102 g 54. How many atoms of hydrogen are present in 12.0 g of water? (Α) 1.1 x 10 23 (Β) 2.0 x 1023 (C) 4.0 x 1023 (D) 8.0 x 1023 (Ε) 4.8 x 1024 55. Which compound contains the highest percent by mass of hydrogen? (Α) HCI

(Β) Η 2 Ο

(D) H2SO4

(Ε) HF

56. Α hydrocarbon (a compound consisting solely of carbon and hydrogen) is found to be 96% carbon by mass. What is the empirical formula for this compound? (Α) C 2 H (Β) CH2 (C) C 3 H (D) CH3 (E) C 4 H 57. An unknown compound contains the elements C, Η, and O. It is known to contain 48% C and 4.0% Η by mass. The molar mass of this compound has been determined in the lab to have a value of 200. The molecular formula for this compound is: SAT CHEMISTRY Dr. D. Bampilis

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(Α) C 2 H 3 O2

(Β) C 4 H6 O 4

(D) C 8 H 3 O 6

(Ε) C8 Η 8 O 6

58. When 7.0 g of ethene (C 2 H4 ) burns in oxygen to give carbon dioxide and water, how many grams of CO 2 are formed? C2 H4 + 3O 2 2CO 2 + 2H 2 O (Α) 9.0 g (Β) 22 g (C) 44 g (D) 82 g (Ε) 180 g 59. Consider the reaction below. What mass of CF4 is formed by the reaction of 8.00 g of methane with an excess of fluorine? CH4(g) + 4F 2( g) CF4(g) + 4HF(g) (Α) 19 g (Β) 22 g (C) 38 g (D) 44 g (Ε) 88 g 60. According to the reaction represented by the unbalanced equation: …SO2(g) +O2(g) …SO3(g) how many moles of SO2(g) are required to react completely with 1 mole of O2(g)? (Α) 0.5 mol (Β) 1 mol (C) 2 mol (D) 3 mol (Ε) 4 mol 61. The combustion of propane, C3 H8(g) , proceeds according to the equation: C3 H8(g) + 5O 2(g) 3CO 2(g) + 4H 2 O (l) How many grams of water will be formed in the complete combustion of 44.0 grams of propane? (Α) 4.50 g (Β) 18.0 g (C) 44.0 g (D) 72.0 g (Ε) 176 g 62. The number of oxygen atoms in 0.50 mole of KHSO4 is: (A) 1.2 x 1023 (B) 2.4 x 1023 (C) 3.0 x 1023 (D) 1.2 x 1024

(Ε) 2.4 x 1024

63. Lithium reacts with water to produce hydrogen gas and lithium hydroxide. What volume of hydrogen collected over water at 22oC and 750 mmHg pressure is produced by the reaction of 0.208 g of Li? [VPH2O = 19.8 mmHg] (A) 367 mL (B) 378 mL (C) 735 mL (D) 755 mL 64. Analysis by mass of a certain compound shows that it contains 14 % hydrogen and 86 % carbon. Which of the following is the most informative statement that can properly be made about the compound on the basis of these data? (A) It is a hydrocarbon. (B) Its empirical formula is CH2. (C) Its molecular formula is C2H4. (D) Its molar mass is 28 g/mol. (E) It contains a triple bond. 65. In a balanced equation, the number of moles of each substance is equal BECAUSE Once the limiting reagent has been consumed, the reaction can no longer continue

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6. Solids, Liquids and Phase Changes Phase Changes In order for a substance to move between the states of matter; for example, to turn from a solid into a liquid, which is called fusion, or from a gas to a liquid (vaporization), energy must be gained or lost. The heat of fusion (symbolized Hfus) of a substance is the amount of energy that must be put into the substance for it to melt. For example, the heat of fusion of water is 6.01 kJ/mol, or in other terms, 80 cal/g. The heat of vaporization, not surprisingly, is the amount of energy needed to cause the transition from liquid to gas, and it is symbolized Hvap. You will not be required to memorize heat of fusion or vaporization values for the exam. Changes in the states of matter are often shown on phase diagrams, and you will probably see at least one of two different types of phase diagrams on the SAT II Chemistry exam. Let’s start with the phase diagram for water. The phase diagram for water is a graph of pressure versus temperature. Each of the lines on the graph represents an equilibrium position, at which the substance is present in two states at once. For example, anywhere along the line that separates ice and water, melting and freezing are occurring simultaneously.

intersection of all three lines is known as the triple point (represented by a dot and a T on the figure). At this point, all three phases of matter are in equilibrium with each other. Point X represents the critical point, and at the critical point and beyond, the substance is forever in the vapor phase. This diagram allows us to explain strange phenomena, such as why water boils at a lower temperature at higher altitudes, for example. At higher altitudes, the air pressure is lower, and this means that water can reach the boiling point at a lower temperature. Interestingly enough, water would boil at room temperature if the pressure was low enough! One final note: If we put a liquid into a closed container, the evaporation of the liquid will cause an initial increase in the total pressure of the system, and then the pressure of the system will become a constant. The value of this final pressure is unique to each liquid and is known as the liquid’s vapor pressure. Water has a relatively low vapor pressure because it takes a lot of energy to break the hydrogen bonds so that molecules enter the gas phase. Water and other liquids that have low vapor pressures are said to be nonvolatile. Substances like rubbing alcohol and gasoline, which have relatively high vapor pressures, are said to be volatile. The

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Example What happens to water when the pressure remains constant at 1 atm but the temperature changes from -10οC to 75οC? Explanation Looking at the phase change diagram for water and following the dashed line at 1 atm, you can see that water would begin as a solid (ice) and melt at 0οC. All of the water would be in liquid form by the time the temperature reached 75οC. The second type of phase change graph you might see on the SAT II Chemistry exam is called a heating curve. This is a graph of the change in temperature of a substance as energy is added in the form of heat. The pressure of the system is assumed to be held constant, at normal pressure (1 atm). As you can see from the graph below, at normal pressure water freezes at 0οC and boils at 100οC.

The plateaus on this diagram represent the points where water is being converted from one phase to another; at these stages the temperature remains constant since all the heat energy added is being used to break the attractions between the water molecules. Specific Heat On the SAT II Chemistry test, you might see a diagram that looks something like this one, and you might come across a question that asks you to calculate the amount of energy needed to take a particular substance through a phase change. This would be one of the most difficult questions on the exam, but you might see something like it, or at least part of it. If you were asked to do this, you would need to use the following equation: energy (in calories) = mCp DT where m = the mass of the substance (in grams) Cp = the specific heat of the substance (in cal/g.οC) DT = the change in temperature of the substance (in either Kelvin or οC, but make sure all your units are compatible!) As you can see, this requires that you know the specific heat of the substance. A substance’s specific heat refers to the heat required to raise the temperature of 1 g of a substance by 1οC. You will not be required to remember any specific heat values for the exam. Work through the example below to get a feel for how to use this equation. Example If you had a 10.0 g piece of ice at -10οC, under constant pressure of 1 atm, how much energy would be needed to melt this ice and raise the temperature to 25.0οC? Explanation First, the temperature of the ice would need to be raised from -10οC to 0οC. This would require the following calculation. The specific heat for ice is 0.485 cal/g.οC. Substituting in the formula energy = mCp DT; energy = (10.0 g) (0.485 cal/g. οC) (10.0 οC) = 48.5 cal So 48.5 calories are needed to raise temperature. SAT CHEMISTRY Dr. D. Bampilis

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Next, we must calculate the heat of fusion of this ice: we must determine how much energy is needed to completely melt the 10 g of it. energy = mH fus energy = (10.0 g) (80 cal/g) = 800 cal So 800 cal of energy are needed to completely melt this sample of ice. Next, we need to see how much energy would be needed to raise the temperature of water from 0ºC to 25ºC. The specific heat for liquid water is 1.00 cal/g.οC. So again use energy = mCp ΔT to get energy = (10.0 g) (1.00 cal/g. οC) (25.0 οC) = 250 cal Finally, add together all of the energies to get the total: 48.5 + 800 + 250 = about 1100 calories are needed to convert the ice to water at these given temperatures.

You have to know: Liquids • Intermolecular interaction • Brownian movement • Viscosity • Surface tension • Capillary action • Boiling point - Vapor pressure • Critical temperature, critical pressure Solids • Crystalline - amorphous • Sublimation - deposition • Melting point • Phase diagram - Triple point • Heating curve Solubility - saturated solution Solid - gas (P, T) Rate of solubility (Pulverizing, Stirring, Heating) Solutions: Dilute - Concentration - Saturated - Unsaturated - Supersaturated Expressions of concentration • % w/w • % w/v • % v/v • Molarity(M) • molality(m)

mol solute Volume solution (L) mol solute 1 kg solvent

Dilution C1·V1 = C2·V2

Exercises For questions 1 - 4: (A) Boyle’s law (B) Charles’ law (C) Avogadro’s law (D) Ideal gas law

(E) Dalton’s law

1. The total pressure of a gaseous mixture is equal to the sum of the partial pressures is: SAT CHEMISTRY Dr. D. Bampilis

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2. Volume is inversely proportional to pressure is: 3. Volume is directly proportional to temperature is: 4. All gases have the same number of moles in the same volume at constant T and P is: For questions 5 -7: (A) Sublimation (B) Condensation (C) Evaporation (D) Deposition 5. Gas

solid is called:

6. Gas

liquid is called:

7. Solid

gas is called:

(E)melting

For questions 8 -10:

(A) AB

(B) BC

(D) DE

(E) EF

8. Which shows melting? 9. Which shows increasing the kinetic energy of a liquid? 10. Which shows boiling? For questions 11 - 19: Statement I BECAUSE Statement II 11. The ideal gas law does not hold under low temperatures and high pressure BECAUSE Interactions between particles cannot be neglected under these conditions 12. CO2 is able to sublimate at atmospheric pressure BECAUSE Its liquid form is impossible to produce 13. When an ideal gas is cooled its volume will increase BECAUSE Temperature and volume are directly proportional 14. According to the KMT, collisions between gas particles and the walls of the container are elastic BECAUSE SAT CHEMISTRY Dr. D. Bampilis

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Gas molecules are considered volume-less particles, with no intermolecular forces, in constant random motion 15. As ice absorbs heat and begins to melt, its temperature remains constant BECAUSE Changes of state bring about changes in a substance’s potential energy, not in its kinetic energy 16. Water boils at a lower temperature at high altitudes compared to low altitudes BECAUSE The vapor pressure of water is lower at higher altitude 17. Decreasing the volume of a system decreases pressure BECAUSE Pressure and volume are inversely related 18. At constant pressure, a certain amount of gas will double in volume as the temperature is halved BECAUSE Temperature and volume are inversely proportional 19. The volume of a gas at 100oC and 600 mmHg will be lower at STP BECAUSE Decreasing temperature and increasing pressure will cause the volume of a gas to decrease 20. What volume would 16 g of molecular oxygen gas occupy at STP? (A) 5.6 L (B) 11.2 L (C) 22.4 L (D) 33.6 L (E) 44.8 L 21. Which of the following is responsible for the abnormally high boiling point of water? (A) Covalent bonding (B) Hydrogen bonding (C) High polarity (D) Large dielectric constant (E) Low molecular weight 22. Which of the following is (are) the weakest attractive forces? (A) Van der Waals (B) Coordinate covalent bonding (C) Covalent bonding (D) Polar covalent bonding (E) Ionic bonding 23. What is the volume at STP of 10 L of gas initially at 546 K, 2 atm? (A) 5 L (B) 10 L (C) 15 L (D) 20 L (E) 25 L 24. If one mole of H2 is compressed from 10 L to 7.5 L at constant temperature, what happens to the gas pressure? (A) It increases by 25% (B) It decreases by 25% (C) It increases by 33% (D) It increases by 50% (E) None of the above

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25. An ideal gas in a closed inflexible container has a pressure of 6 atm and a temperature of 27 deg C. What will be the new pressure at -73 deg C? (A) 2 atm (B) 3 atm (C) 4 atm (D) 8 atm (E) 9 atm For the next few questions, refer to the diagram below, regarding substance Z.

26. Substance Z is at 1 atm and 200 K. If the pressure on substance Z is steadily increased and its temperature is kept constant, what phase change will eventually occur? (A) condensation (B) freezing (C) melting (D) sublimation (E) vaporization 27. The normal boiling point of substance Z is approximately: (A) 100 K (B) 200 K (C) 300 K (D) 400 K

(E) 500 K

28. In what pressure range will the compound sublime? (A) Less than 0.5 atm (B) Between 0.5 and 1.0 (C) Between 1.0 and 2.0 (D) Between 0.5 and 2.0 (E) This compound won’t sublime 29. Crossing line bd is: (A) condensation (B) melting (C) evaporation (D) sublimation (E) boiling 30. Five liters of gas at STP have a mass of 12.5 g. What is the molecular mass of the gas? (A) 12.5 g/mol (B) 25.0 g/mol (C) 47.5 g/mol (D) 56.0 g/mol (E) 125 g/mol 31. Equal molar quantities of hydrogen gas and oxygen gas are present in a closed container at a constant pressure. Which of the following quantities will be the same for the two gases? (A) Partial pressure (B) Partial pressure & average KE (C) Partial pressure & average molecular velocity (D) Average KE & average molecular velocity (E) Partial pressure, average KE, average molecular velocity For the next few questions: A closed 5.0 L vessel contains a sample of neon. The temperature inside the container is 25oC and the pressure is 1.5 atm. 32. Which of the following expressions is equal to the moles of gas in the sample? (A) (1.5 x 5.0) / (0.08 x 25) (B) (0.08 x 250 / (1.5 x 5.0) (C) (1.5 x 25) / (0.08 x 5.0) (D) (0.08 x 298) / (1.5 x 5.0) (E) (1.5 x 5.0) / (0.08 x 298) 33. If the neon gas in the vessel is replaced with an equal molar quantity of helium gas, which will be changed? (A) pressure (B) temperature (C) density (D) pressure & temperature SAT CHEMISTRY Dr. D. Bampilis

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(E) temperature and density 34. The volume was changed while temperature held constant until the pressure was 1.6 atm. Which is equal to the new volume? (A) 5.0 x 1.5 / 1.6 (4.7 L) (B) 5.0 x 1.6 / 1.5 (C) 25 x 1.5 / 1.6 (D) 0.08 x 1.6 / 1.5 (E) 0.08 x 1.5 / 1.6 35. A flask contains three times as many moles of H2 as it does O2. If hydrogen and oxygen are the only gases present, what is the total pressure in the flask if the partial pressure of oxygen is “P”? (A) 4P (B) 3P (C) 4/3P (D) 3/4P (E) 7P 36. The gas in a large cylinder is at a pressure of 3040 torr. Assuming constant temperature and ideal gas behavior, what volume of this gas could you compress into a 100 L box at 8 atm? (A) 20 L (B) 200 L (C) 5000 L (D) 50,000 L (E) 500,000 L 37. Which of the following generalizations CANNOT be made about the phase change of a pure substance from solid to liquid? (A) It involves a change in potential energy (B) It involves no change in temperature (C) It involves a change in kinetic energy (D) It involves a change in entropy (E) It may occur at different temperatures for different compounds 38. If the pressure of a gas sample is doubled at constant temperature, the volume will be: (A) 4 x the original (B) 2 x the original (C) ½ of the original (D) ¼ of the original (E) 1/8 of the original 39. Three canisters, A, B, and C, are all at the same temperature, with volumes of 2.0, 4.0, and 6.0 L, respectively. Canister A contains 0.976 g Ar at 120 torr, Canister B contains 1.37 g N2 at 120 torr, and Canister C is completely empty at the start. Assuming ideality, what would be the pressure in canister C if the contents of A and B are completely transferred to C? (A) 180 torr (B) 330 torr (C) 675 torr (D) 0.25 atm (E) none of the above 40. When a fixed amount of gas has its Kelvin temperature and pressure doubled, the new volume of the gas is: (A) four times greater than its original volume (B) twice its original volume (C) unchanged (D) one half its original volume (E) one fourth its original volume 41. A 600 mL container holds 2 mol O2, 3 mol H2, and 1 mol He. The total pressure within the container is 760 torr. What is the partial pressure of O2? (A) 127 torr (B) 253 torr (C) 380 torr (D) 507 torr (E) 760 torr 42. An ideal gas has a volume of 10 L at 20 deg C and 750 mmHg. Which of the following expressions is needed to determine the volume of the same amount of gas at STP? SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 36

(A) 10 x (750/760) x (0/20) (B) 10 x (750/760) x (293/273) (C) 10 x (760/750) x (0/20) (D) 10 x (760/750) x (273/293) (E) 10 x (750/760) x (273/293) 43. What volume does a sample of 1.50 x 1023 atoms of helium at STP represent? (A) 5.6 L (B) 11.2 L (C) 17.8 L (D) 22.4 L (E) none of the above 44. Which of the following will always decrease the volume of a gas? i. Decrease the pressure with the temperature held constant ii. Increase the pressure with a temperature decrease iii. Increase the temperature with a pressure increase (A) I only (B) II only (C) I and III (D) II and III only

(E) I, II and III

45. A gas has a volume of 10 L at 50 deg C and 200 mmHg. What conversion factor is needed to give a volume at STP? (A) 10 x (0/50) x (200/760) (B) 10 x (0/50) x 760/200) (C) 10 x (273/323) x (200/760) (D) 10 x (273/323) x (760/200) (E) 10 x (323/273) x (760/200) 46. The temperature above which a liquid cannot exist is indicated by: (A) the triple point (B) the critical point (C) the eutectic point (D) the boiling point (E) the sublimation point 47. A change of phase never accompanies: (A) a change in volume (B) a change in pressure (D) a change in density (E) a change in structure

48. The relationship P1V1 = P2V2 is: (A) Boyle’s law (B) Charles’s law (D) the combined gas law (E) the ideal gas law

49. The rate of diffusion of hydrogen gas as compared to that of oxygen gas is: (A) ½ as fast (B) identical (C) twice as fast (D) four times as fast (E) eight times as fast 50. The ratio of the rate of diffusion of oxygen to hydrogen is: (A) 1:2 (B) 1:4 (C) 1:8 (D) 1:16

(E) 1:32

51. Standard conditions using a Kelvin thermometer are: (A) 760 torr, 273 K (B) 760 torr, 273 K, 1 L (C) 760 torr, 0 K (D) 0 torr, 0 K (E) 0 torr, 273 K, 1 L 52. The relation between the pressure and the volume of a gas at constant temperature is given by: (A) Boyle’s law (B) Charles’s law (C) the combined gas law (D) the ideal gas law (E) none of the above SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 37

53. The relation between the absolute temperature and volume of a gas at constant pressure is given by: (A) Boyle’s law (B) Charles’s law (C) the combined gas law (D) the ideal gas law (E) none of the above 54. The relation between the pressure, volume and absolute temperature is given by: (A) Boyle’s law (B) Charles’s law (C) the combined gas law (D) the ideal gas law (E) none of the above 55. At a certain temperature and pressure, ice, water and steam are found to coexist at equilibrium. This pressure and temperature corresponds to: (A) the critical temperature (B) the critical pressure (C) the sublimation point (D) the triple point (E) two of the above 56. How many atoms are present in 22.4 L of O2 at STP? (A) 3 x 1023 (B) 6 x 1023 (C) 9 x 1023 (D) 12 x 1023

(E) 15 x 1023

57. Α gas at STP that contains 6.02 x 1023 atoms and forms diatomic molecules will occupy: (A) 11.2 L (B) 22.4 L (C) 33.6 L (D) 67.2 L (E) 1.06 quarts 58. Inelastic collisions occur in: (A) Real and ideal gases (C) Real gases and fusion reactions

(B) Ideal gases and fusion reactions (D) Real gases (E) Ideal gases

59. The extremely high melting point of diamond (carbon) may be explained by the: (A) network covalent bonds (B) ionic bonds (C) hydrogen bonds (D) van der Waals forces (E) none of the above 60. The phase transition from gas to solid is called: (A) condensation (B) evaporation (C) polymerization

(D) sublimation

61. Which aqueous solution freezes at the lowest temperature? (A) 0.30 m C2H5OH (B) 0.25 m KNO3 (C) 0.20 m CaBr2

(D) 0.10 m FeCl3

62. The value of which concentration unit for a solution changes with temperature? (A) molarity (B) molality (C) mole fraction (D)mass percentage 63. What is the molality of a solution made by dissolving 36.0 g of glucose (C6H12O6, M = 180.2) in 64.0 g of H2O? (A) 0.0533 (B) 0.200 (C) 0.360 (D) 3.12 64. When a nonvolatile solute is dissolved in a volatile solvent, which characteristic is greater for the solution than for the solvent? (A) boiling point (B) freezing point (C) rate of evaporation (D) vapor pressure SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 38

65. Which aqueous solution has the highest osmotic pressure at 25oC? (Assume all ionic compounds ionize completely in solution.) (A) 0.1 M Al2(SO4)3 (B) 0.1 M Na2CO3 (C) 0.2 M KMnO4 (D) 0.3 M C6H12O6 66. A student prepares four 0.10 M solutions, each containing one of the solutes below. Which solution has the lowest freezing point? (A) CaCl2 (B) KOH (C) NaC2H3O2 (D) NH4NO3 67. Ethanol (C2H5OH, M = 46) and methanol (CH3OH, M = 32) form an ideal solution when mixed. What is the vapor pressure of a solution prepared by mixing equal masses of ethanol and methanol? (The vapor pressures of ethanol and methanol are 44.5 mmHg and 88.7 mmHg, respectively.) (A) 133 mmHg (B) 70.6 mmHg (C) 66.6 mmHg (D) 44.5 mmHg 68. A sample of H2 collected over H2O at 23οC and a pressure of 732 mmHg has a volume of 245 mL. What volume would the dry H2 occupy at 0°C and 1 atm pressure? [vp H2O at 23οC = 21 mmHg] (A) 211 mL (B) 218 mL (C) 224 mL (D) 249 mL 69. An ice cube at an unknown temperature is added to 25.0 g of liquid H2O at 40.0οC. The final temperature of the 29.3 g equilibrated mixture is 21.5οC. What was the original temperature of the ice cube? [Cp (J/g.oC) water = 4.184, ice = 2.06, ∆H°fusion = 333 J/g] (A) -6.5οC (B) -13.1οC (C) -35.3οC (D) -56.8οC 70. A sample of a volatile liquid is introduced to an evacuated container with a movable piston. Which change occurs as the piston is raised? (Assume some liquid remains.) I. The fraction of the molecules in the gas phase increases. II. The pressure in the container decreases. (A) I only (B) II only (C) Both I and II (D) Neither I nor II 71. Which point on the phase diagram represents the normal boiling point?

(A) point A

(B) point B

SAT CHEMISTRY Dr. D. Bampilis

(D) point D Σελίδα 39

72. According to the following information, in what physical state(s) does bromine exist at 7.4oC and 400 mmHg? [Triple point -7.3oC, 44 mmHg, Liquid density 3.1 g.cm–3, Solid density 3.4 g.cm–3] (A) solid only (B) liquid only (C) liquid and solid only (D) gas, liquid, and solid 73. The molarity of a solution that is composed of 80.00 g of sodium hydroxide dissolved in 2.0 L of solution is: (Α) 1.0 Μ (Β) 2.0 Μ (C) 4.0 Μ (D) 40.0 Μ (Ε) 160.0 Μ 74. The molarity of a solution obtained when 50.0 mL of 6.0 Μ HCl is diluted to a final volume of 300.0 mL is: (Α) 0.01 Μ (Β) 0.10 Μ (C) 0.20 Μ (D) 0.30 Μ (Ε) 1.0 Μ 75. In the laboratory, a sample of hydrogen is collected by water displacement. The sample of hydrogen has a volume of 25 mL at 24.0 o C and a barometric pressure for the day of 758 mmHg. What is the pressure of the dry gas at this temperature? (The vapor pressure of water at 24.0 o C is 22.4 mmHg.) (Α) 455 mmHg (Β) 470 mmHg (C) 736 mmHg (D) 758 mmHg (Ε) 780 mmHg 76. At the molecular level, the factor that determines whether a substance will be a solid, liquid, or gas is the balance between the: (A) kinetic energies of the molecules and their intermolecular forces. (B) potential energies of the molecules and their intermolecular forces. (C) kinetic energies of the molecules and their intramolecular forces. (D) potential energies of the molecules and their intramolecular forces. 77. The critical temperature of carbon dioxide is 304.3 K. Which statement is true about the behavior of carbon dioxide above this temperature? (A) Solid, liquid and gaseous carbon dioxide are in equilibrium above this temperature. (B) Liquid and gaseous carbon dioxide are in equilibrium above this temperature. (C) Liquid carbon dioxide does not exist above this temperature. (D) Carbon dioxide molecules do not exist above this temperature. 78. The solubility of KClO3 at several temperatures is shown in the accompanying diagram.

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Σελίδα 40

A student mixes 10.0 g of KClO3 with 45.0 g of H2O and stirs it for a long time at 60οC until the solution is completely clear then allows it to cool slowly to 20οC where it remains clear. Which statement about the final clear mixture at 20οC is correct? (A) It is a saturated solution. (B) It is an unsaturated solution and can be made saturated by decreasing the temperature. (C) It is an unsaturated solution and can be made saturated by increasing the temperature. (D) It is a supersaturated solution. 79. Which of the following solutions would probably have the highest boiling point? (Α) 0.100 m ΚΟΗ (Β) 0.100 m Na 2 SO 4 (C) 0.100 m C6H 12O 6 (D) 0.200 m CaCl 2 (Ε) 0.200 m CH3 CH2 OH 80. To determine whether a water solution of Na2S2O3 at room temperature is supersaturated, one can: (A) heat the solution to its boiling point. (B) add water to the solution. (C) add a crystal of Na2S2O3 to the solution. (D) acidify the solution. (E) cool the solution to its freezing point. 81. Which of the following must be measured in order to calculate the molality of a solution? Ι. Mass of the solute. ΙΙ. Mass of the solvent. III. Total volume of the solution. (A) I only (B) I and III only (C) II and III only (D) I and II only (Ε) Ι, ΙΙ, and III 8 2 . Which o f t he fo llo wing so lutes an d stable solutions? SOLUT E I. Ethano l II. Salt III. Oil IV. Oil ( Α ) Ι o nly

( B ) I and II o nly

so lv ents wo ul d be ex pecte d to form SOLVEN T Water Water Vinegar Gaso line

( C ) III o nly

(D) I, II, and III o nly

(E) I, II, and IV o nly 83. Using the sketch of the phase diagram for water given below, determine which of the following statements is incorrect:

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Σελίδα 41

(Α) The triple point is point Α. This is the point at which all three phases are in equilibrium with one another. (Β) The line AB is the line representing the solid-liquid equilibrium line. Anywhere along this line the substance could melt or freeze. (C) The slope of line AB is negative. This slope indicates that the solid is much denser than the liquid. (D) Line AD represents the phase changes of sublimation and deposition. (Ε) Line AC represents where the substance would condense and vaporize. 84. A thermometer is placed in a test tube containing a melted pure substance. As slow cooling occurs, the thermometer is read at regular intervals until well after the sample has solidified. Which of the following types of graphs is obtained by plotting temperature versus time for this experiment?

For questions 85 - 87: Statement I BECAUSE Statement II 85. The solubility of carbon dioxide in a soft drink decreases with a decrease in pressure. BECAUSE The solubility of a gas generally increases with an increase in temperature. 86. Most ionic solids have high melting points. BECAUSE Ionic solids are made up of positive and negative ions held together by electrostatic attractions.

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Σελίδα 42

87. The rate at which sugar dissolves in water increases with stirring. BECAUSE Stirring exposes the surface of a solute crystal to a less concentrated layer of solution. 88. Which statement about the triple point of a substance is correct? (A) The triple point for a substance varies with the pressure. (B) The three phases (solid, liquid, gas) have the same density. (C) The three phases (solid, liquid, gas) are in equilibrium. (D) The three phases (solid, liquid, gas) are indistinguishable in appearance. 89. Under certain conditions CO2 melts rather than sublimes. To which transition in the phase diagram does this change correspond?

(A) A

(B) A

SAT CHEMISTRY Dr. D. Bampilis

(D) C

Σελίδα 43

7. Reaction Types Net Ionic Equations

1. 2. 3. 4. 5. 6. 7.

Net ionic equations are equations that show only the soluble, strong electrolytes reacting (these are represented as ions) and omit the spectator ions, which go through the reaction unchanged. When you encounter net ionic equations on the SAT II Chemistry test, you’ll need to remember the following solubility rules, so memorize them! Also keep in mind that net ionic equations, which are the bare bones of the chemical reaction, usually take place in aqueous environments. Here are those solubility rules: Most alkali metal compounds and NH4+ compounds are soluble. Cl-, Br-, I- compounds are soluble, except when they contain Ag+, Hg22+, or Pb2+. F- compounds are soluble, except when they contain group 2A metals. NO3-, ClO3-, ClO4- and CH3COO- compounds are soluble. SO42- compounds are soluble, except when they include Ca2+, Sr2+, Ba2+, Ag+, Pb2+, or Hg22+. CO32-, PO43-, C2O42-, CrO42-, S2-, OH-, and O2- compounds are insoluble. Group 2A metal oxides are classified as strong bases even though they are not very soluble. Here are some additional rules about common reaction types that you should be familiar with for the exam: If an insoluble precipitate or gas can be formed in a reaction, it probably will be. Oxides (except group 1A) are insoluble, and when reacted with water, they form either acids (nonmetal oxides) or bases (metal oxides). There are six strong acids that completely ionize: HCl, HBr, HI, HNO3, H2SO4, HClO4. All other acids are weak and are written together, as molecules. The strong bases that ionize are oxides and hydroxides of group 1A and 2A metals. All other oxides and hydroxides are considered weak and written together, as molecules. Now try writing some net ionic equations, using the rules above. Example Write the net ionic equation for a mixture of solutions of silver nitrate and lithium bromide. Explanation Ag + + NO3- + Li + + Br This is a double replacement reaction. Both compounds are soluble, so everything ionizes. If anything is formed, it will come from recombining the “inside” two ions with the “outside” two ions to make LiNO3 and AgBr. If either of them is insoluble, a precipitate will be formed, and the ions that react to form it will be in our net ionic equation; the other ions are spectators and should be omitted! As we said, the two possible products are lithium nitrate and silver bromide. Since halides are soluble except those containing silver, mercury, or lead, we have a precipitate of silver bromide, and our net ionic equation looks like this: Ag + + Br AgBr↓ Example Hydrochloric acid and sodium hydroxide are mixed. Write the net ionic equation. Explanation This is a mixture of a strong acid and a strong base, so each ionizes completely. H + + Cl - + Na + + OH The two possible compounds formed are sodium chloride, which is soluble, and water, which is molecular; thus water is the only product in our net ionic equation. H + + OH H2O Example Chlorine gas is bubbled into a solution of potassium iodide; write the net ionic equation.

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Σελίδα 44

Explanation This one is a single replacement, so you need to consider the activity series. Since halogens are involved, you can determine their activity by using the periodic table: Cl is more active than I. Cl 2 + K + + I Remember that halogen is diatomic and that all potassium compounds are soluble. The resulting compound is also soluble, so K+ is a spectator and is left out of the final equation. Cl 2 + I I 2 + Cl -

You have to know: Combination - Synthesis * Zn(s) + S(s) ZnS(s) 2Na(s) + Cl2(g) 2NaCl(s) 2Mg(s) + O2(g) 2MgO(s) H2(g) + Cl2(g) 2HCl(g) H2(g) + 1/2O2(g) H2O(l) 3H2(g) + N2(g) 2NH3(g) C(s) + O2(g) CO2(g) 4Al(s) + 3O2(g) 2Al2O3(s) Decomposition 2HgO(s) 2Hg(s) + O2(g) 2H2O(l) 2H2(g) + O2(g) (electrolysis) MnO

2 2KClO3 2KCl + 3O2(g) Single replacement • Li, Rb, K, Ba, Sr, Ca, Na 2Na(s) + 2H2O(l)cold 2NaOH(aq) + H2(g) Ca(s) + H2SO4(l) CaSO4(s) + H2(g) *(Ba(s) + O2(g) BaO) • Mg, Al, Mn, Zn, Cr, Fe, Cd Mg(s) + H2O(g) MgO(s) + H2(g) Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g) *(2Fe(s) + O2(g) 2FeO(s)) • Co, Ni, Sn, Pb Pb(s) + HCl(aq) PbCl2(aq) + H2(g) * (2Cu(s) + O2(g) 2CuO(s)) • Ag, Pt, Au Double replacement AB + CD AD + CB ΑD or CB: insoluble precipitate gas weak electrolyte (H2O…..) • Soluble:

Na+, K+, NH4+, NO3-, CH3COO-, HCO3-, ClO3• Insoluble: AgCl, HgCl, PCl2 (sol in hot water) BaSO4, BaSO4, PbSO4, Ag2SO4, HgSO4, SrSO4 SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 45

All carbonates, posphates except K+, Na+, NH4+ All sulfides, hydroxides except K+, Na+, NH4+, Ba2+, Ca2+ Gases: HX, H2S, HCN, SO2(H2SO3), CO2(H2CO3), NH3(NH4OH)

Exercises 1 . Indium reacts wit h bro m ine to fo rm In Br 3 . In t he ba lance d equatio n fo r this reactio n, t he co efficient o f t he indium ( ΙΙΙ) bro m i de is: (Α) 1 (Β) 2 (C) 3 (D) 4 (Ε) 6 2 . What is the sum of the coefficients of the following equation when it is balanced? ΑΙ 2 (SO 4 ) 3 + Ca(OH) 2 AI(OH) 3 + CaSO 4 (Α) 5 (Β) 6 (C) 7 (D) 8 (Ε) 9 3 . When the following equation is balanced, what is the sum of the coefficients? ΑΙ 2 (CO 3 ) 3 + M g(OH) 2 AI(OH) 3 + MgCO 3 (Α) 3 (Β) 4 (C) 8 (D) 9 ( Ε ) 10 4. When the equation for the reaction: …BCl3(g) + …H2(g) …HCl(g) + …B(s) is balanced and all coefficients are reduced to lowest whole-number terms, the coefficient for HCl is: (Α) 1 (Β) 2 (C) 3 (D) 4 (Ε) 6 5 . The balanced net ionic equation for the reaction of aluminum sulfate and sodium hydroxide contains which of the following terms? ( Α ) 3 ΑΙ 3 + ( a q ) ( Β ) OH - ( a q ) (C) 3OH- (aq) (D) 2ΑΙ 3+ (aq) (Ε) 2ΑΙ(OH) 3(s) 6. When solutions of phosphoric acid and iron (ΙΙΙ) nitrate react, which of the following terms will be present in the balanced molecular equation? (Α) ΗΝΟ 3(aq) (Β) 3HNO 3(aq) (C) 2FePO 4(s) (D) 3FePO 4(s) (Ε) 2HNO 3(aq) 7. Which solid does not react with a small amount of 3 M HNO3? (A) calcium carbonate (B) manganese (II) sulfide (C) potassium sulfite (D) silver chloride 8. The reaction of silver ion with chloride ion in water solution is a(n): (A) Precipitation (B) Oxidation-reduction (C) Distillation (D) Hydration (E) Condensation 9. The reaction of iron filings with powdered sulfur is a(n): (A) Precipitation (B) Oxidation-reduction (C) Distillation (D) Hydration (E) Condensation 10. A student prepares a 100 mL aqueous solution containing a small amount of (NH4)2SO4 and a second 100 mL solution containing a small amount of Nal, then mixes the two solutions. Which statement describes what happens? (A) Both compounds dissolve and remain in solution when the two solutions are mixed. (B) Both compounds dissolve initially but NH4I precipitates when the solutions are mixed. SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 46

(C) Both compounds dissolve initially but Na2SO4 precipitates when the solutions are mixed. (D) The NaI dissolves but the (NH4)2SO4 does not. There is no change upon mixing. 11. Mixing which pair of 0.10 M solutions produces two precipitates that cannot be separated from one another by filtration? (A) aluminum chloride and copper(II) nitrate (B) strontium bromide and lead(II) acetate (C) magnesium perchlorate and lithium carbonate (D) barium hydroxide and copper(II) sulfate 12. A colored gas is observed with which combination? (A) calcium hydride and water (B) lead metal and nitric acid (C) sodium carbonate and sulfuric acid (D) zinc sulfide and hydrochloric acid 13. Mixing which combination produces a gaseous product? (A) solid ammonium nitrate and solid calcium hydroxide (B) copper metal and 0.10 M hydrochloric acid (C) solutions of barium hydroxide and 0.10 M sulfuric acid (D) solutions of aluminum nitrate and sodium chloride

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8. Thermodynamics You have to know: • Endothermic reaction:

C(s) + O2(g) CO2(g), ΔΗ 0 or C(s) + O2(g) CO2(g) + q • Exothermic reacrion: N2(g) + O2(g)

2NO(g), ΔΗ 0 2NO(g) - q

or N2(g) + O2(g) o f

• Standard enthalpy of formation

o c

• Standard enthalpy of combustion

Calorimetry: q = mcΔΤ q = mL (phase change) Hess’s law rxn

o f

(products)

rxn

o c

(reac tan ts)

o f

(reac tan ts) o c

(products)

Bond dissociation energy (bond length) Bond broken: Energy absorbed Bond formed: Energy evolved o rxn

D broken

Dformed

Entropy – ΔS > 0 • solid → liquid gas •T • mixing different particles •

n (gas) Rxn

ΔG =

ΔΗ -

ΤΔS

Spontaneous

spontaneous at low T

Nonspontaneous

spontaneous at high T

Exercises For questions 1- 3: Refer to the following potential energy diagram & the choices below: -26.4 kcal -67.6 kcal C(s) CO(g) CO2(g) (A) -94.0 kcal

(B) -26.4 kcal

(D) C(s) + ½O2(g)

(Ε) CO2(g)

1. What is the ΔH of the reaction to form CO from C + O2?

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Σελίδα 48

2. What is the ΔH of the reaction to form CO2 from CO + O2? 3. What is the ΔH of the reaction to form CO2 from C + O2? For questions 4 - 6: Refer to the heating curve:

4. In which part of the curve is the state only a solid? 5. In which part is the heat to change the state greatest? 6. In which part is the heat to change the temperature greatest? For questions 7 - 8:

7. Which letter shows the potential energy of the products? 8. Which letter shows the enthalpy change (ΔH) of the reaction? For questions 9 - 13: Refer to the heating curve for H2O below:

9. Where is the temperature of H2O changing at 1oC/g.cal? 10. Which region indicates a solid? 11. Which region indicates a liquid? 12. Which region indicates a gas? 13. Which region indicates a liquid and a gas?

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For questions 14 - 20: Statement I BECAUSE Statement II 14. An exothermic reaction has a positive ΔH BECAUSE Heat is released in an exothermic reaction 15. A calorimeter can be used to measure the amount of heat lost or absorbed in a process BECAUSE Calorimeters can be used to measure heat lost or gained by a system and its surroundings 16. The freezing of water is an exothermic process BECAUSE Energy is released when covalent bonds are formed 17. An increase in entropy leads to a decrease in randomness BECAUSE The low energy state of ordered crystals has high entropy 18. An exothermic reaction has a positive ΔH value BECAUSE Heat must be added to an exothermic reaction for the reaction to occur 19. Covalent bonds must be broken for a liquid to boil BECAUSE Heat is released when a liquid changes into a gas 20. The temperature of a substance always increases as heat energy is added to it BECAUSE The average kinetic energy of the particles in a system increases with an increase in temperature 21. How much heat is given off when 8 g of hydrogen reacts in: 2H2 + O2 2H2O? ΔH = -115.60 kcal (A) -57.8 kcal (B) -115.6 kcal (C) -173.4 kcal (D) -231.2 kcal (Ε) -462.4 kcal 22. A reaction that absorbs heat is: (A) endothermic (B) an equilibrium process (D) non-spontaneous (Ε) exothermic

23. The change in heat energy for a reaction is best expressed as a change in: (A) Enthalpy (H) (B) Absolute temperature (T) (C) Specific heat (c) (D) Entropy (S) (Ε) Kinetic energy (KE) 24. When 1 mole of sulfur burns to form SO2, 1300 calories are released. When 1 mole of sulfur burns to form SO3, 3600 calories are released. What is ΔH when 1 mole of SO2 burns to form SO3? SAT CHEMISTRY Dr. D. Bampilis

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(A) 3900 cal

(B) -1950 cal

(D) -500 cal

(Ε) -2300 cal

25. When the temperature of a 20 gram sample of water is increased from 10 oC to 30oC, the heat absorbed by the water is (A) 600 cal (B) 30 cal (C) 400 cal (D) 20 cal (Ε) 200 cal 26. How many g of CH4 produce 425.6 kcal in: CH4 + 2O2 (A) 8 g (B) 16 g (C) 24 g (D) 32 g

CO2 + 2H2O + 212.8 kcal (Ε) 64 g

27. 10 g of liquid at 300 K are heated to 350 K. The liquid absorbs 6 kcal. What is the specific heat of the liquid (in cal/g.oC)? (A) 6 (B) 120 (C) 12 (D) 600 (Ε) 60 28. CH4(g) + 2O2(g) CO2(g) + 2H2O(g) + 800 kJ. If a mole of O2(g) is consumed in the reaction, what energy is produced? (A) 200 kJ (B) 400 kJ (C) 800 kJ (D) 1200 kJ (Ε) 1600 kJ 29. What is ΔHrxn for the decomposition of 1 mole of NaClO3? ΔHf = -85.7 kcal/mol for NaClO3(s) and ΔHf = -98.2 kcal/mol for NaCl(s) (A) -183.9 kcal (B) -91.9 kcal (C) +45.3 kcal (D) +22.5 kcal (Ε) -12.5 kcal 30. What is the heat of combustion of one mole of C2H4? Compound ΔHf (kcal/mol) H2O(g) -57.8 C2H4(g) 12.5 CO2(g) -94.1 (A) +316.3 kcal (B) -12.5 kcal (C) -291.3 kcal (D) -316.3 kcal (Ε) -57.8 kcal 31. Given 2Na(s) + Cl2(g) 2NaCl(s) + 822 kJ, how much heat is released if 0.5 mol of sodium reacts completely with chlorine? (A) 205.5 kJ (Β) 411 kJ (C) 822 kJ (D) 1644 kJ (Ε) 3288 kJ 32. Calculate the approximate amount of heat necessary to raise the temperature of 50.0 grams of liquid water from 10.0 o C to 30.0 o C. (The specific heat of water liquid is 4.18 J/g. o C.) (Α) 20 J (Β) 80 J (C) 100 J (D) 200 J (Ε) 4,180 J 33. In neutralizing 500 mL of 1.0 Μ HCI with 500 mL of 1.0 Μ ΝaΟΗ, the temperature of the solution rises 5.0 ο C. Given that the density of the solution is 1.0 g/mL and the specific heat of the solution is 4.184 J/g. o C, calculate the approximate energy released from this experiment. (Α) 20 J (Β) 1000 J (C) 4200 J (D) 2.1 x 104 J (Ε) 1.0 x 10 4 kJ 34. Use the bond energies given below to estimate the enthalpy, ΔΗ, for the following reaction: C 2 H 4 +CI 2 CIH2C-CH2 CI

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Bond energies kJ/mol C-C 347 C=C 612 C-Cl 341 C-H 414 Cl-Cl 243 (Α) ΔH = -800 kJ (Β) ΔH = -680 kJ (C) ΔH = -150 kJ (D) ΔH = +150 kJ (Ε) ΔH = +200 kJ 35. Based on the relationship of entropy to the degree of disorder of a system, which response includes all the occurrences listed that represent a decrease in entropy? I. The freezing of water. II. The vaporization of water. III. Sublimation (vaporization) of dry ice, solid CO2. IV. The extraction of Mg and pure water from seawater. (A) I and II (B) II and IV (C) I and IV (D) III (E) II and III 36. Consider the reaction below. When a 45.00 gram sample of ethanol is burned with excess oxygen, about how much energy is released as heat? C 2 H5 OH( l) + 3O 2(g) 2CO 2(g) + 3Η 2 O (l) ΔΗ = -1.40 x 10 3 kJ ( Α ) 0.995 kJ (Β) 5.1x10 2 kJ (C) 1.40 x 10 3 kJ (D) 2.80 x 103 kJ (Ε) 5000 kJ 37. Consider the following hypothetical reaction (at 375Κ). The standard free energies in kJ/mol are given below each substance. What is the value of the Gibb's free energy for the reaction at this temperature? Is the reaction spontaneous? 2Α Β + C ΔG O =? -20.0 150.0 -350.0 (Α) -220; yes (Β) -180; no (C) -160; yes (D) +180; no (Ε) -160; no 38. How much heat is required to convert 5.0 g of ice at –10.0oC to liquid water at 15.0oC? (Assume heat capacities are independent of temperature.) Enthalpy of fusion: 6.00 kJ.mol–1, Specific heat capacity of ice:37.8 J.mol–1.oC–1, Specific heat capacity of water:76.0 J.mol–1.oC–1 (A) 4.2 x 102 J (B) 2.1 x 103 J (C) 9.3 x 103 J (D) 3.8 x 104 J 39. Each of two solutions is mixed separately, and both solutions are found to be the same temperature. The two solutions are mixed, and a thermometer shows that the mixture's temperature has decreased in temperature. Which of the following statements is true? (Α) The chemical reaction is exothermic. (Β) The chemical reaction is absorbing energy. (C) The chemical reaction is releasing energy. (D) The energy released could be found by multiplying the temperatures together. (Ε) The energy absorbed by the solution is equal to the difference in temperature of the solutions. 40. Spontaneous reactions are driven by: (A) Low enthalpy values and high entropy values. (B) Low enthalpy values and low entropy values. SAT CHEMISTRY Dr. D. Bampilis

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(C) High enthalpy values and low entropy values. (D) High enthalpy values and high entropy values. (Ε) High temperatures and low pressures. 41. If an exothermic process is spontaneous, which of the following statements must be true? (Α) ΔG must be positive. (Β) ΔS must be positive. (C) ΔS must be negative. (D) The temperature must be over 500Κ. (Ε) ΔG must be negative. For question 42: Statement I BECAUSE Statement II 42. The reaction presented by the equation: 2H2 ( g ) + O2( g) 2 Η 2 O ( I ) , Δ Η o R X N = -5 72 kJ is exothermic. BECAUSE The total enthalpy of the products in this reaction is less than that of the reactants.

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9. Kinetics You have to know: Factors affecting reaction rates • nature of reactants • surface area • concentration • temperature • catalyst Collision theory Activation energy Reaction rate law - Mechanism Zero order reactions First order reactions

Rate [A]t

[ ] k[A] t [A] e kt

When t

1/ 2

1 : [ ] 2

[ ]e

1/ 2

1 2

1/ 2

ln 2 k

0.693 k

Exercises For questions 1 - 4: (A) Positive ΔH (B) Negative ΔH (C) Positive ΔG (D) Negative ΔG (E) Positive ΔS 1. Which describes an endothermic reaction? 2. Which describes a spontaneous reaction? 3. Which describes a nonspontaneous reaction? 4. Which is multiplied by temperature in the equation which calculates free energy? For questions 5 - 7: (A) An increase in the reaction concentration (C) A decrease in pressure (D) catalysis

(B) An increase in temperature (E) pH

5. Which increases effective collisions without increasing average kinetic energy? 6. Which decreases the activation energy? 7. Which increases average kinetic energy? For questions 8 - 10: SAT CHEMISTRY Dr. D. Bampilis

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(A) Enthalpy change (D) Activation energy

(Β) Entropy change (C) Gibbs free energy change (Ε) specific heat capacity

8. Which is the amount of energy that must be added to raise the temperature of 1 gram of a substance 1oC? 9. Which value indicates the spontaneity of a reaction? 10. Which value indicates whether a reaction is endothermic or exothermic? For questions 11 - 13: (Α) 0 (Β) 1 (C) 2

(D) 3

(E) 4

Based on the following reaction and rate data: A(g) + B(g) products Experiment [A] mol/L [B] mol/L Initial rate (M/s) 1 0.060 0.010 0.040 2 0.030 0.010 0.040 3 0.030 0.020 0.080 11. What is the order of the reaction with respect to A? 12. What is the order of the reaction with respect to B? 13. What is the rate constant? For questions 14 - 26: Statement I BECAUSE Statement II 14. An increase in entropy leads to a decrease in randomness BECAUSE The low energy state of ordered crystals has high entropy 15. In a second order reaction doubling [A] quadruples the rate BECAUSE The rate equation is r = k[A]2 for such a reaction 16. Catalysts decrease the rate of a chemical reaction BECAUSE Catalysts decrease activation energy 17. An exothermic reaction has a positive ΔH value BECAUSE Heat must be added to the reaction for the reaction to occur 18. The entropy of a solid decreases when it is dissolved BECAUSE It becomes less ordered SAT CHEMISTRY Dr. D. Bampilis

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19. If ΔH is positive and ΔS is positive, ΔG must be positive BECAUSE ΔG = ΔH – TΔS 20. A catalyst increases the rate of reaction BECAUSE It raises the energy of the products 21. Increasing the temperature increases the reaction rate BECAUSE At high temperatures, molecules or atoms tend to be further apart 22. Activation energy of a reaction is decreased by a catalyst BECAUSE It is not used up in the process 23. The reaction of BaCl2 and Na2SO4 doesn’t go to completion BECAUSE The compound barium sulfate is formed 24. Catalysts speed up or slow down a reaction BECAUSE Catalysts change the energy released from a reaction 25. When salt dissolves in water, ΔS for the process is positive BECAUSE Aqueous ions have greater entropy than ions in a solid 26. The addition of a catalyst will decrease the ΔH for a reaction BECAUSE Catalysts provide alternate reaction paths with lower activation energy 27. For the reaction: A + B C, determine the order of the reaction with respect to B from the information given below: [A]o [B]o Initial rate (M/s) 1.00 1.00 2.0 1.00 2.00 8.1 2.00 2.00 15.9 (A) Zero order (Β) First order (C) Second order (D) Third order (Ε) Fourth order 28. Which of the following is NOT a true statement about entropy? (A) Entropy is a measure of the randomness in a system (Β) The entropy of an amorphous solid is greater than that of a crystalline solid (C) The entropy of a spontaneous reaction cannot decrease SAT CHEMISTRY Dr. D. Bampilis

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(D) The entropy of an isolated system will spontaneously increase or remain constant (Ε) The entropy of a liquid is greater than that of a solid 29. A catalyst: (A) changes ΔG for an equation (Β) acts by increasing the rate of the forward reaction more than the reverse reaction (C) raises the equilibrium constant of a system (D) may have molecular weight as low as 1 or higher than 200,000 (Ε) does not react chemically during the course of a reaction 30. For which of the following is there an increase in entropy? (A) Na(s) + H2O(l) NaOH(aq) + H2(g) (Β) I2(g) I2(s) (C) H2SO4(aq) + Ba(OH)2(aq) BaSO4(s) + H2O(l) (D) H2(g) + ½ O2(g) (Ε) None of the above

H2O(l)

31. Which of the following conditions guarantee a spontaneous reaction? (A) positive ΔH, positive ΔS (Β) positive ΔH, negative ΔS (C) negative ΔH, negative ΔS (D) negative ΔH, positive ΔS (Ε) none of the above 32. A reaction that occurs only when heat is added is best described as: (A) Exothermic (Β) Endothermic c. An equilibrium process (D) Spontaneous (Ε) Non-spontaneous 33. The important considerations in deciding if a reaction will be spontaneous are: (A) stability & state of reactants (Β) energy gained & heat evolved (C) exothermic energy & randomness of the products (D) endothermic energy & randomness of the products (Ε) endothermic energy & structure of the products 34. In a multistep chemical process, the rate-limiting step is the step in the reaction with the: (A) Highest activation energy & fastest reaction rate (Β) Highest activation energy & slowest reaction rate (C) Lowest activation energy & fastest reaction rate (D) Lowest activation energy & slowest reaction rate (Ε) Greatest concentration of the reactants and products 35. Which of the following reactions shows a decrease in entropy? (A) C(s) + 2H2(g) CH4(g) (Β) H2O(g) H2(g) + ½ O2(g) (C) 2NI3(s) N2(g) + 3I2(g) (D) 2O3(g) 3O2(g) (Ε) none of the above For questions 36-37 use the following diagram:

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36. Which letter corresponds to the activation energy of the reaction? (Α) Α

(Β) B

(D) Y

(E) X

37. Which letter corresponds to the change in energy for the overall reaction? (Α) Α (Β) B (C) C (D) Υ (Ε) Χ For question 38: Statement Ι BECAUSE Statement ΙI 38. The reaction shown above is exothermic. BECAUSE Energy difference B is greater than energy difference A. 39. What is the first-order rate constant for a reaction that is 36.5% complete in 0.0200 seconds? (A) 50.4 s–1 (B) 27.7 s–1 (C) 22.7 s–1 (D) 9.86 s–1 40. When a sample of an ideal gas is heated from 25oC to 50oC the average kinetic energy of the molecules increases. Which ratio gives the correct relationship between the average kinetic energies at the higher temperature to the lower temperature? (A) 2 : 1

(B)

(D)

41. The catalyzed pathway in a reaction mechanism has a _______activation energy and thus causes a _____ reaction rate. (Α) higher, lower (Β) higher, higher (C) lower, higher (D) lower, steady (Ε) higher, steady 42. Which of the following statements about catalysts is true? (A) They increase the value of the equilibrium constant. (B) They increase the amount of product present at equilibrium. (C) They increase the concentration of the reactants. (D) They are permanently altered as the reaction proceeds. (E) They reduce the activation energy of the reaction. 43. Which of the following statements best describes the condition(s) needed for a successful formation of a product in a chemical reaction? SAT CHEMISTRY Dr. D. Bampilis

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(Α) The collision must involve a sufficient amount of energy, provided from the motion of the particles, to overcome the activation energy. (Β) The relative orientation of the particles has little or no effect on the formation of the product. (C) The relative orientation of the particles has an effect only if the kinetic energy of the particles is below some minimum value. (D) The relative orientation of the particles must allow for formation of the new bonds in the product. (Ε) The energy of the incoming particles must be above a certain minimum value and the relative orientation of the particles must allow for formation of new bonds in the product.

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10. Equilibrium You have to know: Equilibrium constant Kc - Kp (solid)

k1 k2

k: Rate constants

Le Chatelier principle (C, P, Θ) H2O(l)

H2O(g) o

Kp – ΔG: ΔG = - RTlnK small K ΔGo positive K = e-ΔG/RT

Exercises 1. What is correct about the signs and magnitudes of the free energy, ∆G˚, and the equilibrium constant, K, for a thermodynamically spontaneous reaction under standard conditions? (A) ∆Go < 0, K < 0 (B) ∆Go = 0, K > 0 (C) ∆Go < 0, K = 0 (D) ∆Go < 0, K > 0 2. For the process, CH3OH(l) CH3OH(g) ∆G° = 4.30 kJ/mol at 25°C. What is the vapor pressure of CH3OH(l) at 25°C in mmHg? (A) 0.176 mmHg (B) 14.0 mmHg (C) 134 mmHg (D) 759 mmHg 3. For Br2, ∆H˚vap = 31 kJ . mol–1. If S˚ values for Br2(g) and Br2(l) are 245 J . mol–1 . K–1 and 153 J . mol-1 . K–1 respectively, what is the normal boiling point for Br2(l)? (A) 340 K (B) 200 K (C) 130 K (D) 70 K 4. Which range includes the value of the equilibrium constant, Keq, for a system with ∆G˚ << 0? (A) –1 < Keq < 0 (B) 0 < Keq < 1 (C) Keq < –1 (D) 1 < Keq 5. For the reaction at 25oC, C2H4(g) + H2(g) What is ∆G˚ for this reaction in kJ.mol-1? (A) –0.436 (B) –3.71

C2H6(g) Kp = 5.67×107 (C) –19.2

(D) –44.2

6. The gaseous compound NOBr decomposes according to the equation: NOBr(g) NO(g) + 1/2 Br2(g) At 350 K the equilibrium constant, Kp, is 0.15. What is the value of ∆G˚? (A) –5.5 × 103 J/mol (B) –2.4 × 103 J/mol (C) 2.4 × 103 J/mol (D) 5.5 × 103 J/mol

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7. Tungsten is obtained commercially by the reduction of WO3 with H2 according to the equation: WO3(s) + 3H2(g) W(s) + 3H2O(g) The following data related to this reaction at 25oC are available. WO3(s) H2O(g) ∆H˚ kJ/mol –840.3 –763.5 ∆G˚ kJ/mol –763.5 –228.5 The temperature at which this reaction is at equilibrium at 1 atm is closest to which of the following? (A) 124 K (B) 213 K (C) 928 K (D) 2810 K 8 . Consider the system below at equilibrium. Which of the following changes will shift the equilibrium to the right? N2(g) + 3H2 (g ) 2 NH3 ( g ) + 9 2 .9 4 kJ Ι. Increasi ng t he tem perat ure. ΙΙ. D ecreasing t he tem perature. ΙΙΙ. I ncreasing t he pres sure o n t he system . ( A ) I o nly ( B ) II o nly ( C ) III o nly (D) Ι and III ( Ε ) ΙΙ and III 9 . The value of the equilibrium constant, Κ, is dependent on: Ι. The tem perature o f t he sy stem . ΙΙ. The co nce ntratio n o f t he reactants. ΙΙΙ. The co ncentratio n o f t he pro ducts. IV. T he nature o f t he reactants an d pro ducts . ( Α ) I, II ( Β ) ΙΙ, ΙΙΙ ( C ) III, IV (D) Ι and IV

( Ε ) Ι, ΙΙ, and IV

1 0 . If at a given tem perature t he e quilib rium co nstant fo r t he reactio n: H2 ( g ) +Cl 2 ( g ) HCl ( g ) (Α)

2HCl ( g ) is Κ p , the e quilibrium co nstant fo r t he reactio n: 1/ 2 H2 ( g ) +1 / 2 Cl 2 ( g ) can be re presented as: (Β)

(C)

(D)

(E)

11. Write the equilibrium expression for the following reaction: 2Α (g) + B (g) 3C(s) + D (g) (A) [A] 2 [B][D] ( B )

(C)

(D)

(Ε)

12. For the hypothetical reaction: A + B C + D, the equilibrium constant, K, is less than 1.0 at 25oC and decreases by 35% on changing the temperature to 45oC. What must be true according to this information? (A) The ∆H˚ for the reaction is negative. (B) The ∆S˚ for the reaction is positive. (C) The ∆G˚ for the reaction at 25o is negative. (D) The ∆G˚ for the reaction at 45o is zero.

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13. Which of the following events is least likely to occur with an increase in temperature for the reaction given? (With ΔΗ = -45.9 kJ/mol.) Ν 2 + 3Η 2 2ΝΗ 3 (Α) The gas particles will move more quickly. (Β) The reaction will produce more ammonia in a shorter time. (C) The reaction will reverse and ammonia will decompose. (D) The entropy of the system will increase. (Ε) The equilibrium constant will become smaller. For questions 14 - 16:Statement Ι BECAUSE Statement ΙI 14. Α system is at equilibrium when the rate of the forward reaction is equal to the rate of the reverse reactions. BECAUSE At equilibrium, the concentration of the products is equal to that of the reactants. 15. In the system N2(g) + O2(g) ⇌ 2NO(g) decreasing the pressure will not cause a shift in position of the equilibrium BECAUSE There is no net change in the number of moles of gas from one side of the reaction to another 16. When the temperature of a reaction at equilibrium is increased, the equilibrium will shift to favor the endothermic direction BECAUSE Endothermic reactions involve heat acting as a reactant and Le Chatelier’s principle states that an equilibrium shift will occur to offset temperature changes 17. BaCl2 dissociates in water to give one Ba2+ ion and two Cl- ions. If concentrated HCl is added to this solution: (A) [Ba2+] increases (Β) [OH-] increases (C) [Ba2+] remains constant (D) [H+] decreases (Ε)the number of moles of undissociated BaCl2 increases 18. Consider: H2(g) + Br2(g) ⇌ 2HBr(g) The concentrations of H2, Br2 and HBr are 0.05 M , 0.03 M, and 500.0 M. The equilibrium constant for this reaction at 400oC is 2.5 x 103. Is this system at equilibrium? (A) Yes, the system is at equilibrium (Β) No, the reaction must shift to the right in order to reach equilibrium (C) No, the reaction must shift to the left in order to reach equilibrium (D) It cannot be determined (Ε)The reaction will never be at equilibrium 19. A chemist interested in the reactivity of iodine concentrates his study on the decomposition of gaseous hydrogen iodide: 2HI(g) ⇌ H2(g) + I2(g). What is the equilibrium expression for this reaction? (A) [H2]2[I2] (Β) [H2] (C) [H2][I2]/[HI]2 (D) [H2][I2]2 (Ε) [H2]2[I2]2 SAT CHEMISTRY Dr. D. Bampilis

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20. An equilibrium expression may be forced to completion by: (A) adding a catalyst (Β) increasing the pressure (C) increasing the temperature (D) removing the products from the reaction mixture as they are formed (Ε) decreasing the reactant concentration 21. The Haber process is used for producing ammonia from nitrogen and hydrogen. This reaction could be forced to produce more ammonia by: (A) increasing the reaction pressure (Β) decreasing the reaction pressure (C) adding a catalyst (D) both (B) and (C) (Ε)none of the above 22. An increase in pressure will change the equilibrium by: (A) shifting to the side where a smaller volume results (Β) shifting to the side where a larger volume results (C) favoring the endothermic reaction (D) favoring the exothermic reaction (Ε)None of the above 23. Which statement is true for a liquid/gas mixture at equilibrium? (A) The equilibrium constant is dependent on temperature (Β) The amount of the gas present at equilibrium is independent of pressure (C) All interchange between the liquid and gas phases has ceased (D) All of the above (Ε)None of the above 24. The equilibrium expression, K = [CO2] represents the reaction: (A) C(s) + O2(g) ⇌ CO2(g) (Β) CO(g) + ½ O2(g) ⇌ CO2(g) (C) CaCO3(s) ⇌ CaO(s) + CO2(g) (D) CO2(g) ⇌ C(s) + O2(g) (Ε)CaO(s) + CO2(g) ⇌ CaCO3(s) 25. In this equilibrium reaction: A + B ⇌ AB + heat, in a closed container, how could the forward reaction rate be increased? i. Increasing [AB] ii. Increasing [A] iii. Removing some of AB (A) i only (Β) iii only (C) i and iii only (D) ii and iii only (Ε)i, ii and iii 26. An increase in pressure in the reaction 2HI(g) ⇌ H2(g) + I2(g) would: (A) produce more I-(aq) (Β) produce more H2 (C) not affect the system (D) drive it to the right (Ε)drive it to the left

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11.Solutions

Solution

Suspension Colloid Cloudy, heterogeneous, at Cloudy but uniform and least two substances homogeneous visible

Appearance

Clear, transparent and homogeneous

Particle Size

molecule in size

Effect of Light Tyndall Effect

none -- light passes through, particles do variable not reflect light

Effect of Sedimentation

none

larger than 10,000 Angstroms

particles will eventually settle out

10-1000 Angstroms

light is dispersed by colloidal particles

none

Colligative Properties Properties of solutions that depend on the number of solute particles present per solvent molecule are called colligative properties. The concentration of solute in a solution can affect various physical properties of the solvent including its freezing point, boiling point, and vapor pressure. For the SAT II you will only need to be familiar with the first two. Freezing Point Depression The freezing point of a substance is defined as the temperature at which the vapor pressure of the solid and the liquid states of that substance are equal. If the vapor pressure of the liquid is lowered, the freezing point decreases. Why is a solution’s freezing point depressed below that of a pure solvent? The answer lies in the fact that molecules cluster in order to freeze. They must be attracted to one another and have a spot in which to cluster; if they act as a solvent, solute molecules get in the way and prevent them from clustering tightly together. The more ions in solution, the greater the effect on the freezing point. We can calculate the effect of these solute particles by using the following formula: DTf = Kf x m solute x i where: DTf = the change in freezing point Kf = molal freezing point depression constant for the substance (for water = 1.86oC/m) m = molality of the solution i = number of ions in solution (this is equal to 1 for covalent compounds and is equal to the number of ions in solution for ionic compounds) Boiling Point Elevation As you learned earlier in this chapter, the boiling point of a substance is the temperature at which the vapor pressure equals atmospheric pressure. Because vapor pressure is lowered by the addition of a nonvolatile solute, the boiling point is increased. Why? Since the solute particles get in the way of the solvent particles trying to escape the substance as they move around faster, it will take more energy for the vapor pressure to reach atmospheric pressure, and thus the boiling point increases. We can calculate the change in boiling point in a way that’s similar to how we calculate the change in freezing point: SAT CHEMISTRY Dr. D. Bampilis

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DTb = Kb x m solute x i where: Kb = molal boiling point elevation constant (for water = 0.51oC/m) Now try a problem that deals with freezing point depression and boiling point elevation.

Example Calculate the freezing point and boiling point of a solution of 100 g of ethylene glycol (C2H6O2) in 900 g of water.

Explanation Calculate molality:

Freezing point depression = (m)(Kf)(i) Tf = (1.79)(1.86)(1) = 3.33 o C Freezing point = 0 o C - 3.33 o C = -3.33 o C Boiling point elevation = (m)(Km)(i) Tb = (1.79)(0.51)(1) = 0.91 o C Boiling point = 100 o C + 0.91 o C = 100.91 o C

You have to know: Colligative properties •

P Po

N N n

Po P Po

P: V.P. of solution

n N n

N: moles of solvent

Po: V.P. of solvent n: moles of solute

• Δθb = kb × (molality) kb = 0.52o C/1m • Δθf = kb × (molality) kb = -1.86oC/1m molality of ionic solutions! • Π V = n RT For questions 1 - 2: (A) Molarity (Β) Formality

(D) Molality

(Ε)Normality

1. Which is defined as moles of solute per kilogram of solvent? 2. Which is defined as moles of solute per liter of solution? For questions 3 - 6: (A) Solute (Β) Solvent

(D) Aqueous solution

(Ε)Solvation

3. Which is present in a lesser amount in a solution? 4. Which describes a species in which the solvent is water? 5. Which is present in greater quantity in solution? SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 65

6. Which is the interaction between the solute and the solvent molecules? For questions 7 - 9: (A) Dilute (Β) Concentrated (D) Saturated (Ε)Supersaturated

7. Which is the condition, unrelated to quantities, that indicates that the rate going into solution is equal to the rate coming out of solution? 8. Which is the condition that exists when a water solution that has been at equilibrium is heated to a higher temperature with a higher solubility, but no additional solute is added? 9. Which is the descriptive term that indicates there is a large quantity of solute, compared with the amount of solvent, in a solution? For questions 10 - 11: (A) i, ii and iii (Β) i and ii

(D) i only

(Ε)iii only

10. The ionization of salts in water is useful in explaining: i. Their unusually large solubility in water ii. Their electrical conductivity in solution iii. The lowered freezing points and increased boiling points of their solutions 11. Which of the following describe water? i. Solvent for polar solutes ii. Polar molecule

iii. Good conductor of electricity

For questions 12 - 17: Statement I BECAUSE Statement II 12. When a solute is added to pure water, the vapor pressure of the water will decrease BECAUSE All solutes dissociate into positive and negative ions 13. NaCl(aq) is an electrolyte BECAUSE It forms ions in solution 14. A salt dissolved in an organic solvent will be a good electrical conductor BECAUSE Salts will not dissolve appreciably in an organic solvent 15. A super saturated solution of glucose in boiling water crystallizes as it cools BECAUSE The solubility increases as the temperature decreases 16. Salt dissolved in water lowers the freezing point SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 66

BECAUSE The change in freezing point is given by: ΔTf = iKfm 17. Sodium chloride forms aqueous solution of ions BECAUSE The sodium has a +1 charge and the chloride has a -1 charge and they are hydrated by the water molecules 18. What is the molarity if a 500 mL solution contains 20 g of CaBr2? (A) 0.1 M (Β) 0.2 M (C) 0.5 M (D) 1 M (Ε)5 M 19. How many moles of sulfate ions are in 200 mL of a 2 M sodium sulfate solution? (A) 0.2 mol (Β) 0.4 mol (C) 0.6 mol (D) 0.8 mol (Ε)1.0 mol 20. A 0.5 m solution could be prepared by dissolving 20 g NaOH in: (A) 0.5 L water (Β) 0.5 kg water (C) 1 L water (D) 1 kg water (Ε)2 L water 21. What volume of water would be needed to dilute 50 mL of a 3 M HCl solution to 1 M? (A) 25 mL (Β) 50 mL (C) 75 mL (D) 100 mL (Ε)150 mL 22. About how many grams of sodium chloride would be dissolved in water to form a 0.5 M solution in 500 mL solution? (A) 7 (Β) 29 (C) 14.5 (D) 58 (Ε)112 23. A one liter solution of 2 M NaOH can be prepared with: (A) 20 g NaOH (Β) 40 g NaOH (C) 60 g NaOH (D) 80 g NaOH (Ε)100 g NaOH 24. What is the molarity of a 10 mL solution in which 3.7 g KCl are dissolved? (A) 0.05 M (Β) 0.1 M (C) 1 M (D) 5 M (Ε)10 M 25. A solution of 10 M NaOH was used to prepare 2 L of 0.5 M NaOH. How many mL of the original NaOH solution are needed? (A) 10 mL (Β) 100 mL (C) 1000 mL (D) 200 mL (Ε)2000 mL 26. A 1 molal solution of NaCl results when 58.5 g of sodium chloride is dissolved in: (A) One liter of water (Β) 100 mL of water (C) one kilogram of water (D) 100 g of water (Ε)one cubic meter of water 27. A small crystal of NaCl is added to a sodium chloride solution resulting in the precipitation of more than 1 gram of sodium chloride. This solution had been: (A) Unsaturated (Β) Saturated (C) Supersaturated (D) Dilute (Ε)Concentrated 28. Sodium chloride would be most soluble in: (A) Ether (Β) Benzene (C) Water (D) Carbon tetrachloride SAT CHEMISTRY Dr. D. Bampilis

(Ε)Gasoline Σελίδα 67

29. Which of the following would produce a highly conductive aqueous solution? (A) Cyclohexane (Β) Hydrochloric acid (C) Benzene (D) Sucrose (Ε)Acetic acid 30. A 10% solution of HNO3 would be produced by dissolving 63 g HNO3 in how many mL water? (A) 100 (Β) 300 (C) 567 (D) 630 (Ε)1000 31. What is the boiling point of an aqueous solution containing 117 g NaCl in 1000 g H2O? Kb = 0.52oC kg/mol (A) 98.96oC (Β) 99.48oC (C) 100.52oC (D) 101. 04oC (Ε)102.08oC 32. Which of the following sequences lists the relative sizes of particles in a water mixture from smallest to largest? (A) Solutions, suspensions, colloids (Β) Solutions, colloids, suspensions (C) Colloids, solutions, suspensions (D) Colloids, suspensions, solutions (Ε)Suspensions, colloids, solutions 33. A compound which, when dissolved in water, barely conducts electrical current can probably be: (A) A strong electrolyte (Β) An ionic salt (C) A strong acid d. A strong base (Ε)None of the above 34. How many grams of HCl must be added to 500 mL water to produce a solution that freezes at -1.86oC? (molal freezing constant = 1.86oC kg/mol) (A) 4.6 (Β) 9.1 (C) 18.3 (D) 36.5 (Ε)73.0 35. A solution can be both: (A) Dilute and concentrated (Β) Saturated and dilute (C) Saturated and unsaturated (D) Supersaturated and saturated

(Ε)None of these

36. The solubility of a solute indicates: (A) The temperature of the solution (Β) The quantity of solvent (C) The quantity of solute (D) The nature of the solute & solvent (Ε)All of these 37. A 10% solution of NaCl means that in 100 g of solution there is: (A) 5.85 g NaCl (Β) 58.5 g NaCl (C) 10 g NaCl

(D) 94 g NaCl

38. The molarity of a solution made by placing 98 g H2SO4 in sufficient water to make 500 mL of solution is: (A) 0.5 (Β) 1 (C) 2 (D) 3

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Σελίδα 68

39. If 684 g sucrose (MM = 342 g) is dissolved in 2000 g H2O (essentially 2 L), what will be the freezing point of this solution? (A) -0.51oC (B) -1.86oC (C) -3.72oC (D) -6.58oC 40. Ten grams of sodium hydroxide dissolved in 1 L of water makes a solution that is: (A) 0.25 M (B) 0.5 M (C) 1 M (D) 1.5 M (Ε) 4 M 41. How much water, in liters, must be added to 0.5 L of 6 M HCl to make it 2 M? (A) 0.33 (B) 0.5 (C) 1 (D) 1.5 (Ε) 2 42. How many grams of NaOH are needed to make 100 g of a 5 % solution? (A) 2 (B) 5 (C) 20 (D) 40 (Ε) 95 43. What is the boiling point of water at the top of a mountain? (A) 100oC (B) > 100oC since the pressure is less than at ground level (C) < 100oC since the pressure is less than at ground level (D) > 100oC since the pressure is greater than at ground level (Ε)< 100oC since the pressure is greater than at ground level 44. 1 mol NaCl in 1000 g H2O will change the boiling point of water to: (A) 100.51oC (B) 101.04oC (C) 101.53oC (D) 101.86oC (Ε)103.62oC 45. To what volume, in mL, must 50.0 mL of 3.50 M H2SO4 be diluted in order to make 2 M H2SO4? (A) 25 (B) 60.1 (C) 87.5 (D) 93.2 (Ε)101 46. What is the molar mass of a non-ionizing solid if 10 g of this solid, dissolved in 100 g of water, formed a solution that froze at -1.21oC? (A) 0.65 g (B) 6.5 g (C) 130 g (D) 154 g (Ε)265 g 47. What is the melting point of 0.2 L water containing 6.20 g C2H6O2? (A) -1.86oC (B) -0.93oC (C) 0oC (D) 0.93oC (Ε)1.86oC 48. Which of the following are TRUE? i. Adding a solute raises the vapor pressure & boiling point. ii. The change in boiling & freezing point depends on molality. iii. The number of solute particles in a solvent is an important factor in determining the boiling point elevation. (A) i only (B) ii only (C) i and ii (D) ii and iii (Ε)i, ii and iii

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Σελίδα 69

12. Acids, Bases, Salts You have to know: Arrhenius - Brönsted / Lowry - Lewis Properties of acids • Water solutions conduct electricity • React with Me to liberate H2(g) • Change the color of indicators (phenolphthalein: colorless – litmus: red) • React with bases (neutralization) • React with carbonates to release CO2(g) • Strong: HNO3, HCl, H2SO4, HI, HBr, HClO4 Properties of bases • Water solutions conduct electricity • Change the color of indicators (phenolphthalein: pink – litmus: blue) • React with acids (neutralization) • React with fats to make soap • Aqueous solutions feel slippery and stronger bases are very caustic to the skin • Strong KOH, NaOH, Ba(OH)2, Sr(OH)2, Ca(OH)2 Brönsted - Lowry • Conjugate Acid - Base • Stronger acid weaker its conjugate base • HClO4(aq) + H2O(l) H3O+(aq) + ClO4-(aq) stronger acid stronger base weaker acid weaker base General conclusion: is that proton transfer reactions favor the production of the weaker acid and base Amphoteric species (H2O, HA-) Lewis Acid ionization constants

[H 3O ][A ] [HA]

[

] 55.6

mol L

Ka - Strength of acids Base ionization constants kb Ionization constant of water Kw = [H3O+].[OH-] At 25oC Kw = 10-14 pH + pOH = 14 pH = - log[H3O+] = - log[H+] pOH = - log[OH-] pH scale Common ion effect - (Le Chatelier’s principle) Buffer solutions Titration Indicators SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 70

Hydrolysis of salts Solubility products Ksp - Qsp Ksp - S AgCl(s) AgSO4(s)

Ag+(aq) + Cl-(aq) Ksp = s2

2Ag+(aq) + SO42-(aq) Ksp = 4s3

Normality Salts formation • Neutralization • Single replacement • Direct combination Fe(s) + S(s) • Double replacement • MexOy + AmezOw Acid rain •

SO2(g) O2(g)

Ksp 4

FeS(s)

SO3(g)

SO3(g) H 2O(l)

• 2NO 2(g)

s = K sp

H 2SO4(aq)

H 2 O (l)

HNO 2(aq)

HNO3(aq)

Exercises For questions 1 - 4: (A) HBr(aq) (B) NH3(aq)

(D) HF(aq)

(Ε)H2CO3(aq)

1. A strip of litmus paper will appear blue in 2. At 25oC, it has a pH > 7 3. It is essentially a non-electrolyte 4. Its aqueous ionization goes virtually to completion For questions 5 - 8: (A) a Bronsted acid (Ε) a buffer

(B) a Bronsted base

(D) a weak base

5. It’s a solution made by the combination of a weak acid and the salt of its conjugate base 6. It always dissociates completely in aqueous solution 7. It has a very high Ka 8. It accepts a proton For questions 9 - 12: (A) a strong acid

(B) a strong base

SAT CHEMISTRY Dr. D. Bampilis

(D) a weak base Σελίδα 71

(Ε)a salt (made from an acid and a base) 9. NH3 is 10. Cl- is 11. NaHCO3 is 12. NaOH is For questions 13 - 18: (A) an acid (B) a base (Ε)an amphoteric substance

(D) a basic salt

13. Amino acids are an example of 14. Ammonia is an example of a 15. Ammonium sulfate is an example of 16. Aluminum chloride is an example of 17. The product of a group IA element and water is an example of a 18. Bicarbonate ion is an example of a For questions 19 - 32: Statement I BECAUSE Statement II 19. The reaction of zinc with hydrochloric acid goes to completion in an open container Because Hydrogen gas is evolved from the reaction of zinc and hydrochloric acid. 20. A 0.2 M solution of carbonic acid is a weaker conductor of electricity than a 0.2 M solution of HBr BECAUSE In solutions with the same concentration of solute molecules, H2CO3 is less dissociated than HBr 21. An aqueous solution of HI is considered to be a Bronsted - Lowry base. BECAUSE HI(aq) can accept an H+ ion from another species. 22. If an acid is added to pure water, it increases the water’s pH. BECAUSE Adding an acid to water raises the hydrogen ion concentration in the water.

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Σελίδα 72

23. Hydrofluoric acid etches glass. BECAUSE It is a strong acid. 24. Acetic acid is a strong acid. BECAUSE Acetic acid ionizes completely in solution. 25. NH3 is a Lewis base. BECAUSE Ammonia can accept a proton. 26. A 1 N (“normal”) solution of H2SO4 is the same as a 1M (“molar”) solution of H2SO4. BECAUSE Molarity refers to the moles of solute per liter of solution, whereas normality refers to the molarity of hydrogen ions. 27. The pH of 0.01 M HCl(aq) is 2. BECAUSE HCl is essentially an ionic species, completely dissociating so that [H+] = [HCl]. 28. A solution with a pH of 12 has a higher concentration of hydroxide ions than a solution with a pH of 10 BECAUSE At 25oC, pH + pOH = 14. 29. A basic solution has more hydrogen ions than an acidic solution. BECAUSE At 25oC, the product of [H+] x [OH-] = 10-14. 30. Water makes a good buffer BECAUSE A good buffer will resist changes in pH 31. When volumes of 1.0 M HCl and 1.0 M NaOH are mixed, the product mixture is theoretically safe to drink. BECAUSE The acid and the base form a neutral salt 32. If an acid is added to water with original pH of 7, the concentration of hydroxide ions will increase. BECAUSE The product of hydroxide ions and hydrogen ions is equal to 1.0 x 10-14 in all aqueous solutions at 25oC.

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Σελίδα 73

33. In HNO3(aq) + OH-(aq) ⇌ H2O(l) + NO3-(aq), which species is the conjugate acid? (A) HNO3(aq) (B) OH-(aq) (C) H2O(l) (D) NO3-(aq) (Ε)There is no conjugate acid 34. Which is true regarding an aqueous solution of H3PO4 at 25oC? (A) It has a very large acid ionization constant (B) It has a bitter taste (C) The concentration of [OH-] > 1.0 x 10-7 M (D) It is a weak electrolyte (Ε)It can be formed by the reaction of a metal oxide and water 35. In NH3(aq) + H2CO3(aq) ⇌ NH4+(aq) + HCO3-(aq), NH4+(aq) acts as a(n): (A) indicator (B) hydrate (C) acid (D) base (Ε)salt 36. Which of the following are true regarding the aqueous dissociation of HCN, Ka = 4.9 x 10-10 at 25oC? i. At equilibrium, [H+] = [CN-] ii. At equilibrium, [H+] = [HCN] iii. HCN is a strong acid (A) i only (B) ii only (C) i and ii only (D) ii and iii only (Ε)i, ii and iii 37. The reaction of zinc metal and HCl produces which of the following? i. H2(g) ii. Cl2(g) iii. ZnCl2(aq) (A) ii only (B) iii only (C) i and ii only (D) i and iii only (Ε)i, ii and iii 38. Which characteristic is associated with Lewis bases? (A) React with metal to produce hydrogen gas (B) Donate an unshared electron pair (C) Always contain the hydroxide ion in its structure (D) Taste sour (Ε)Formed by the reaction of a nonmetal oxide and water 39. Which of the following is a poor electrolyte? (A) A hydrochloric acid solution (B) A sodium hydroxide solution (C) A vinegar solution (D) A sodium chloride solution (Ε)Molten sodium chloride 40. A compound that dissolves in water which barely conducts electrical current can probably be (A) A strong electrolyte (B) An ionic salt (C) A strong acid (D) A strong base (Ε)None of the above 41. Which of the following acids is capable of dissolving gold? (A) Hydrochloric (B) Nitric (C) Sulfuric (D) A combination of A and B (Ε)A combination of A and C 42. A stock solution of 10 M NaOH was used to prepare 2 L of 0.5 M NaOH. How many milliliters of sodium hydroxide stock solution were used? (A) 10 mL (B) 100 mL (C) 1000 mL (D) 200 mL (Ε)2000 mL

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Σελίδα 74

43. What is the hydroxide ion concentration in a solution with a pH of 5? (A) 10-3 (B) 10-5 (C) 10-7 (D) 10-9 (Ε)10-11 44. What is the H3O+ concentration of a 0.100 M acetic acid solution (Ka = 1.8 x 10-5)? (A) 1.8 x 10-5 (B) 1.8 x 10-4 (C) 1.3 x 10-2 (D) 1.3 x 10-3 (Ε)0.9 x 10-3 45. What is the pH of a solution with a hydroxide ion concentration of 0.00001 M? (A) -5 (B) -1 (C) 5 (D) 9 (Ε)14 46. A titration experiment is conducted in which 15 mL of a 0.015 M Ba(OH)2 solution is added to 30 mL of an HCl solution of unknown concentration and titration is complete. What is the approximate concentration of the HCl solution? (A) 0.015 M (B) 0.03 M (C) 1.5 M (D) 2.5 M (Ε)3.0 M 47. An aqueous solution with pH 5 at 25oC has a hydroxide ion concentration of (A) 1 x 10-11 M (B) 1 x 10-9 M (C) 1 x 10-7 M (D) 1 x 10-5 M (Ε)1 x 10-3 M 48. What is the pOH of a solution with [H+] = 0.001 M (A) -3 (B) 1 (C) 3 (D) 11

(Ε)14

49. Which of the following can be used to prepare hydrogen gas in the laboratory? (A) Mercuric oxide (B) Acid plus zinc (C) Potassium chlorate (D) Carbon disulfide (Ε)Benzene 50. What is the solubility in pure water of Ba(IO3)2 in moles per liter at 25oC? [Ksp (25oC) = 6.0 x 10–10] (A) 1.2 x 10–5 (B) 1.7 x 10–5 (C) 5.3 x 10–4 (D) 8.4 x 10–4 51. A 500 mL saturated solution of MgCO3 (Mr = 84) is reduced to 120 mL by evaporation. What mass of solid MgCO3 is formed? [Ksp = 4.0 × 10–5] (A) 0.0013 g (B) 0.064 g (C) 0.20 g (D) 0.27 g 52. For the dissolution of Ag2SO4, ∆H˚ = 17.8 kJ . mol–1 and ∆S˚ = –34.9 J . mol–1 . K–1 at 25oC. What is the value of the Ksp for Ag2SO4 at this temperature? (A) 5.0 × 10–2 (B) 7.6 × 10–4 (C) 5.3 × 10–4 (D) 1.1 × 10–5 53. The solubility of AgBrO3 in aqueous solution depends on the presence of other substances in solution. Relative to its solubility in H2O the solubility of AgBrO3 is higher in 0.10 M ____ and lower in 0.10 M ____. (A) NH3, KBrO3 (B) KBrO3, NH3 (C) HNO3, NH3 (D) NH3, HNO3 54. Equal volumes of 1 x 10–4M solutions of Cd2+and CO32– ions are mixed in one flask and equal volumes of 1 x 10–4M solutions of Ag+ and CrO42– ions are mixed in a second. Which substances precipitate?

SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 75

(A) CdCO3 only

Formula CdCO3 Ag2CrO4 –12 Ksp 5.2 x 10 1.1 x 10–12 (B) Ag2CrO4 only (C) Both

(D) Neither

55. In a solution with the standard molarity, which of the following has the greatest [Η 3 O + ]? (Α) HCN (Β) ΗΝΟ 3 (C) H2 O (D) ΟΗ (Ε) CH3 OH 56. Which of the following would produce a basic aqueous solution? (Α) SO 2 (Β) KCI (C) CO 2 (D) NH4 CI

(Ε) Na 2 O

57. Equimolar solutions of which of the following would produce the most acidic solution? (Α) Η 3 ΡΟ 4 (Β) HClO (C) HClO 2 (D) HClO 3 (Ε) HClO 4 58. All of the following can act as Bronsted - Lowry acids (proton donors) in aqueous solution EXCEPT: (Α) HI (Β) NH 4 + (C) HCO 3 (D) H 2 S (Ε) NH 3 59. When HS- acts as a Bronsted base, which of the following is formed? (Α) S 2(Β) H + (C) H 2 S (D) H 2 S 2

(Ε) H 3 S +

60. A solution is made by adding 5.6 g of KOH (molar mass: 56 g/mol) to enough water to make 1.0 L of solution. What is the approximate pH of the resulting solution? (Α) 1 (Β) 3 (C) 7 (D) 9 (Ε) 13 61. The hydrogen ion concentration of a solution prepared by diluting 50 ml of 0.10 M HNO3(aq) with water to 500 ml of solution is: (Α) 0.0010 M (Β) 0.0050 M (C) 0.010 M (D) 0.050 M (Ε) 1.0 M 62. Which of the following solutions is weakly acidic? (Α) 0.1 M HCl (Β) 0.1 M NaCl (C) 0.1 M HC 2 H 3 O 2

(D) 0.1 M CH 3 OH (Ε) 0.1 M KOH

63. Which of the following solutions has the highest pH? (Α) 0.1 M HCl (Β) 0.1 M NaCl (C) 0.1 M HC 2 H 3 O 2 (D) 0.1 M CH 3 OH (Ε) 0.1 M KOH 64. Which of the following solutions reacts with an equal volume of 0.05 M Ba(OH)2 to form a solution with pH = 7? (Α) 0.1 M HCl (Β) 0.1 M NaCl (C) 0.1 M HC 2 H 3 O 2 (D) 0.1 M CH 3 OH (Ε) 0.1 M KOH 65. For which salt is the molar solubility, s, equal to 4×10–6M? (A) AgC2H3O2 Ksp = 2×10–3 (B) TlBr Ksp = 4×10–6 (C) MnCO3 Ksp = 2×10–11 (D) Zn(OH)2 Ksp = 3×10–17

SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 76

66. Which silver compound is the most soluble in water? (A) AgCl Ksp = 1.8×10-10 (B) Ag2CO3 Ksp = 8.5×10-12 (D) Ag3PO4 Ksp = 8.9×10-17

67. Which of the following substances would dissociate completely when placed into excess amounts of distilled water? (Α) C2H5OH (Β) HC2H3O2 (C) LiNO3 (D) Mg(OH)2 (Ε) All of these will dissociate completely in water 68. Which of the following is not true for a solution at 25 o C that has a hydroxide concentration of 1.0 x 10 -8 Μ? (Α) Κ w =1x10 -14 . (Β) The solution is acidic. (C) The solution is basic. (D) The [Η + ] is 1 x 10 -8 Μ. (Ε) The pΟΗ equals 6.0.

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Σελίδα 77

13. Redox and Electrochemistry You have to know: Oxidation: increase in oxidation number Reduction: decrease in oxidation number Oxidation: Loss of Electrons Reduction: Gain of Electrons Rules for Assigning an Oxidation State Oxidizing - Reducing agent Balancing redox equations - The oxidation state method - Using half equations

(e / H / H 2 O

H 2O / H / e )

Electrochemistry cells (voltaic - electrolytic - difference from IB) Anode: oxidation occurs Cathode: reduction occurs Voltaic/galvanic cells (electricity is produced) e.g. a cell Cu/Zn Anode (-)

Cathode (+) Cu

Zn 2

(Oxidation) + 0.76 V (Reduction) + 0.34 V Eo cell + 1.10 V > 0 (reactions occurs)

Standard reduction potentials (Eo) (SOS! The Eo are not multiplied) Eo high: oxidizing agent Eo low: reducing agent Nernst equation: E

RT ln K eq nF

2.30 RT log Q (1) nF nFE

RT ln K

nFE

F: faraday is the charge of one mole of electrons = 96500 C n: number of moles of electrons transferred (1) E

0.0591 log Q (T n

298 K)

Electrolytic cell (electrical energy required to induce reaction) (-) Reduction - Cathode: most oxidizing cation (less reactive metal) or H2O (+) Oxidation - Anode: most reducing anion (less reactive non metal) or H2O Molten KCl, Aqueous NaCl, H2O (H+), electroplating

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Σελίδα 78

Exercises For 1 - 5: (A) 0 (B) -1

1. The oxidation number of Na in NaCl 2. The oxidation number of Cl in Cl2 3. The oxidation number of S in Na2S 4. The charge of calcium in CaCl2 5. The charge of chlorine in KCl For 6 - 8: (A) Zn(s)

(B) Cu2+(aq)

(D) Cu(s)

(Ε)H2O

6. Acts as the anode 7. Acts as the cathode 8. Is reduced For 9 - 13: (A) group IA

(B) group IIA

(D) group VIA

(Ε)group VIIA

9. _________3(PO4)2 10. _________2O2 (oxidation state of oxygen is -1) 11. Cu_____2 12. Good reducing agents 13. Group represented by the Lewis dot structure below X For 14 - 15: (A) 1 (B) 2

(D) 4

(Ε)5

14. When the following equation: HMnO4 + H2SO3 coefficient, in the lowest whole number, of H2SO3 is 15. When the following equation Br2 + SO2 + H2O in the lowest whole number, of HBr is SAT CHEMISTRY Dr. D. Bampilis

MnSO4 + H2O + H2SO4 is balanced, the

H2SO4 + HBr is balanced, the coefficient,

Σελίδα 79

Q Statement I BECAUSE Statement II 16. Cu2+ ion needs to be oxidized to form Cu metal BECAUSE Oxidation is the gain of electrons 17. The anions migrate to the cathode in an electrochemical reaction BECAUSE Positively charged ions are attracted to the negatively charged electrode 18. The alkali metals are strong oxidizing agents BECAUSE The one electron in their valence shell is easily lost 19. The standard reduction potential for Ag+ + e- Ag is half that of 2Ag+ + 2eBECAUSE Standard potential is dependent on the number of electrons transferred

2Ag

20. Chloride ions, Cl-, can be oxidized to produce chlorine gas BECAUSE Two chloride ions give up an electron to form Cl2 21. The oxidation state of Cr in Al2(Cr2O7)3 is +3 BECAUSE As a neutral compound, the sum of the oxidation numbers of all the atoms must equal zero 22. The electrolysis of potassium iodide, KI, produces electrical energy BECAUSE Electrolytic cells convert chemical energy into electrical energy 23. An ionic solid is a good conductor of electricity BECAUSE An ionic solid is composed of positive and negative ions joined together by electrostatic forces 24. Elemental sodium is a good reducing agent BECAUSE An atom of elemental sodium gives up its valence electron readily 25. What’s the potential of the reaction below given the half-reaction potentials: 2Fe2+ + Cl2 2Fe3+ + 2Cl-? Fe3+ + e- Fe2+; E = 0.77 V Cl2 + 2e- 2Cl-; E = 1.36 V (A) 0.18 V (B) 0.59 V (C) 1.05 V (D) 2.13 V (Ε)2.90 V

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26. For Cu(s) + NO3-(aq) + H+(aq) Cu2+(aq) + NO2(g) + H2O(l), when the equation is balanced what is the coefficient of H+? (A) 1 (B) 2 (C) 3 (D) 4 (Ε)5 27. For Cu(s) + NO3-(aq) +H+(aq) Cu2+(aq) + NO2(g) + H2O(l), which of the following takes place? (A) Cu(s) is oxidized (B) H+(aq) is oxidized (C) Cu(s) is reduced (D) H+(aq) is reduced (Ε)NO3- is oxidized 28. The standard reduction potential of Cu2+(aq) is +0.34 V. What is the oxidation potential of Cu(s)? (A) +0.68 V (B) +0.34 V (C) -0.34 V (D) -0.68 V 29. If the following reactions are used to make a galvanic cell, which species will be reduced and which species will be oxidized? F2 + 2e- 2F-(aq); E = +2.87 V Ca2+ + 2e- Ca(s); E = -2.76 V (A) F- will be oxidized and Ca2+ will be reduced (B) Ca2+ will be oxidized and F2 will be reduced (C) Ca(s) will be oxidized and F2 will be reduced (D) F2 will be oxidized and Ca(s) will be reduced 30. What is the oxidation number of Mn in KMnO4? (A) -7 (B) -3 (C) 0 (D) +3 (Ε)+7 31. Which of the following is true of an electrolytic cell? (A) An electric current causes an otherwise non-spontaneous chemical reaction to occur. (B) Reduction occurs at the anode (C) A spontaneous electrochemical reaction produces an electric current (D) The electrode to which the electrons flow is where oxidation occurs (Ε)None of the above 32. What is the sum of the coefficients of the products for the following reaction? K2Cr2O7 + HCl KCl + CrCl3 + H2O + Cl2 (A) 10 (B) 12 (C) 13 (D) 14 (Ε)15 33. The oxidation number of sulfur in NaHSO4? (A) 0 (B) +2 (C) -2 (D) +4 (Ε)+6 34. How many moles of electrons are required to reduce 103.6 g of lead from Pb2+ to the metal? (A) 0.5 mole (B) 1 mole (C) 2 moles (D) 4 moles (Ε)8 moles 35. The order of decreasing strength as reducing agents is: (A) Na, Mg, Fe, Ag, Cu (B) Mg, Na, Fe, Cu, Ag (C) Ag, Cu, Fe, Mg, Na SAT CHEMISTRY Dr. D. Bampilis

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(D) Na, Fe, Mg, Cu, Ag (Ε)Na, Mg, Fe, Cu, Ag 36. Electrolysis of a dilute solution of aqueous sodium chloride results in the cathode product (A) Sodium (B) Hydrogen (C) Chlorine d. Oxygen (Ε)Peroxide 37. For the following reactions: Zn Zn2+ + 2e-; E = +0.76 V Au Au3+ + 3e-; E= -1.42 V If gold foil is placed in a solution containing Zn2+, the reaction potential would be: (A) -1.34 V (B) -2.18 V (C) -0.66 V (D) +2.18 V (Ε)+1.34 V 38. In the electrolysis of molten copper chloride, the substance liberated at the anode is (A) Copper (B) Chlorine (C) Hydrogen (D) Copper chloride (Ε)None of the above 39. The deposition of 1.0 g of which element from its molten chloride requires the shortest time at a current of 1 A? (A) Na (B) Mg (C) Al (D) Ba 40. The cell: Al(s) | Al3+(aq) 0.001 M) | | Cu2+(aq) 0.10 M) | Cu(s) has a standard cell potential, E˚ = 2.00 V. What is the cell potential for this cell at the concentrations given? (A) 2.07 V (B) 2.03 V (C) 1.97 V (D) 1.94 V 41. The equilibrium constant, K, is 2.0 × 1019 for the cell Ni(s) / Ni2+(aq) // Hg22+(aq) / Hg(l) The value of E˚ at 25oC for this cell is closest to: (A) –1.14V (B) –0.57V (C) 0.57V (D) 1.14V 42. Chromium metal can be produced by the electrolysis of molten CrO3. What current in amperes operating for 100 minutes is needed to produce 104 grams of this metal? (A) 193 (B) 96.5 (C) 64.3 (D) 32.2 43. For the reaction: C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(l), ∆G˚ = –2.108 × 103 kJ.mol–1. What is the value of the standard electrode potential, E˚ for a fuel cell based on this reaction? (A) 1.09 V (B) 2.18 V (C) 4.37 V (D) 21.8 V 44. What is the [Fe2+] in a cell at 25oC for which E = –0.458 V vs a standard hydrogen electrode? Fe2+(aq) + 2e– Fe(s) E° = –0.440 V (A) 0.246 M (B) 0.496 M (C) 2.01 M (D) 4.06 M 45. What must be the pH in the hydrogen compartment of the cell designated below if the cell voltage is 0.70 V? Zn(s) / Zn2+(aq) // H+(aq) / H2(g) E˚ = 0.76 V (Assume that both the [Zn2+] and the H2(g) pressure are at standard values and T = 25oC.) (A) 0.51 (B) 1.01 (C) 2.50 (D) 3.21 SAT CHEMISTRY Dr. D. Bampilis

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46. Ethanol reacts with dichromate ions in acid solution according to the equation: C2H5OH(l) + Cr2O72–(aq) + H+(aq) CO2(g) + Cr3+(aq) + H2O(l) What is the coefficient for H+(aq) when this equation is balanced with the smallest whole number coefficients? (Α) 10 (B) 12 (C) 14 (D) 16 (E)48 47. Ag+(aq) + e– Ag(s) E° = 0.80 V 2+ – Mg (aq) + 2e Mg(s) E° = –2.73 V Use the equations above to calculate the value of ∆G° (in kJ/mol) for the reaction: Mg(s) + 2Ag+(aq) Mg2+(aq) + 2Ag(s) (A) 681 (B) 341 (C) –341 (D) –681 48. Consider the following reactions: X(NO3)2 + Y X + Y(NO3)2 X(NO3)2 + Z X + Z(NO3)2 Y(NO3)2 + Z No reaction What is the correct order of increasing activity for the metals X, Y, Z? (A) X < Y < Z (B) X < Z < Y (C) Z < Y < X (D) Z < X < Y 49. Given these standard reduction potentials: Ag+(aq) + e– Ag(s) E˚ = 0.80 V 2+ – Pb (aq) + 2e Pb(s) E˚ = –0.13 V what is the free energy change (in kJ.mol-1) for the reaction: Pb(s) + 2Ag+(aq) Pb2+(aq) + 2Ag(s)? (A) –180 (B) –90 (C) 90

(D) 180

50. For the voltaic cell represented: Ni(s) Ni2+(aq) Ag+(aq) Ag(s) which change will increase the cell potential? (A) increasing the [Ag+] (B) increasing the [Ni2+] (C) adding Ni(s)` (D) removing Ag(s) 51. In which of the following substances would nitrogen have the highest oxidation number? (A)NO (B)N 2 O (C)NO 2 (D)N 2 O 4 (E)NO 3 52. When the equation: …Cu2+(aq) + …I-(aq) …CuI(s) + …I2(s) is balanced and all coefficients are reduced to lowest whole-number terms, the coefficient for I-(aq) is: (Α) 1 (Β) 2 (C) 3 (D) 4 (Ε) 5 53. Consider a voltaic cell in which the reaction below occurs in two half-cells connected by a salt bridge and an external circuit. 2Cr(s) + 3Sn2+(aq) 3Sn(s) + 2Cr3+(aq) E˚ = 0.603 V Which change will cause the voltage to increase? (A) Increasing the amount of Sn(s) in its half-cell (B) Increasing the amount of Cr(s) in its half-cell (C) Diluting the solution in the anode compartment SAT CHEMISTRY Dr. D. Bampilis

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(D) Diluting the solution in the cathode compartment 54. For the cell: Zn(s) + 2H+(aq) Zn2+(aq) + H2(g) E˚ = 0.76 V Which change will increase the voltage of the cell? (A) Increasing the size of the Zn electrode. (B) Increasing the [Zn2+]. (C) Increasing the [H+]. (D) Increasing the pressure of the H2(g). 55. Sn(s) | Sn2+(aq) || Cu2+(aq) | Cu(s) For the voltaic cell represented above, which change will increase the voltage? (A) Increasing the size of the Sn electrode (B) Increasing the size of the Cu electrode 2+ (C) Increasing the [Sn ] (D) Increasing the [Cu2+] 56. Given the following chemical reaction for the formation of lithium oxide, which of the following statements is true? 4Li (s) + O 2(g) 2Li 2 O (s) (Α) Lithium metal is the oxidizing agent. (Β) Oxygen gas is the reducing agent. (C) Lithium is oxidized. (D) Oxygen is oxidized. (Ε) Oxygen loses two electrons to become a -2 ion. 57. Which of the following half-cell reactions describes what is happening at the anode in the diagram?

(Α) Ζn Ζn 2+ + 2e (D) SΟ 4 S + 2O2 + 6e -

(Β) Η 2 2Η + + 2e (Ε) 2Η + + 2e Η2

Cl 2 + 2e -

58. Voltaic cells harness the energy of redox reactions. BECAUSE In a voltaic cell, electron flow occurs through the salt bridge. 59. Electrolytic cells require the input of energy. BECAUSE Electrolytic cells have just one container, while voltaic cells have two. 60. Zinc metal will reduce Cu+2 in solution. BECAUSE Zinc is a more active metal than copper is.

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14. Organic Chemistry You have to know: Alkanes CnH2n+2 sp3 - tetrahedral • combustion • CH4 + Cl2/Br2 • nonpolar Alkyl groups R - or CnH2n+1 Cycloalkanes CnH2n Aromatics

resonance structure

ortho - meta - para Alkenes CnH2n sp2 - 120o C n H 2n Br2 sp - 180o

Alkynes CnH2n-2

Alcohols Cn H 2n 1OH Ethers Cn H 2n

O Cm H 2m

Aldehydes Cn H 2n

C H

Ketones Cn H 2n

C Cm H 2m

Carboxylic acids Cn H 2n

C OH

Esters Cn H 2n

C O Cm H 2m

Amines Cn H 2n 1 NH 2 Amides Cn H 2n

C NH 2 ||

a Aminoacids Cn H 2n

C H C OH |

NH2

Exercises Q Statement I BECAUSE Statement II 1. Carbon is a nonmetal BECAUSE Carbon atoms can bond with each other 2. The hybrid orbital form of carbon in acetylene is believed to be the sp form BECAUSE It is a linear compound with a triple bond between carbons 3. Normal butyl alcohol and 2-butanol are isomers SAT CHEMISTRY Dr. D. Bampilis

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BECAUSE Isomers vary in the number of neutrons in the nucleus of the atom 4. The reaction of CaCO3 and HCl goes to completion BECAUSE Reactions that form a precipitate go to completion 5. The alkanes are considered a homologous series BECAUSE Homologous series have the same functional group but differ in the formula by the addition of a fixed group of atoms 6. Benzene is a poor electrolyte in water solution BECAUSE It does not ionize 7. Benzene does not have true single and double bonds between its carbon atoms in the ring BECAUSE It is composed of delocalized pi electrons in the ring giving rise to resonance structures 8. Long chain hydrocarbons are insoluble in water BECAUSE “like dissolves like” and water contains oxygen and no carbon and long chain hydrocarbons contain carbon, but no oxygen 9. Ethylene (C2H4) has a higher carbon-carbon bond energy than acetylene BECAUSE Ethylene contains a double bond and acetylene has only a single bond between the carbons 10. Benzene (C6H6) can be drawn as a series of resonance structures BECAUSE Its bonds are a hybrids of single and double bond character 11. Which of the following has the strongest carbon-carbon bond? (Α) C2H2 (B) C2H4 (C) C2H6 (D) C2H8

(Ε) C2H10

12. Which of the following statements is true of ethene? (Α) Both carbon atoms are sp2 hybridized and the molecule is planar (B) Both carbon atoms are sp2 hybridized and all bond angles are approximately 109.5 o (C) One carbon atom is sp hybridized while the other is sp2 (D) Both carbon atoms are sp3 hybridized and all bond angles are approximately 109.5 o (Ε) Both carbon atoms are sp hybridized and the molecule is planar 13. Which of the following is the formula for a non-cyclic, saturated hydrocarbon? (Α) C7H12 (B) C7H14 (C) C7H16 (D) C7H18 (Ε) C7H20 SAT CHEMISTRY Dr. D. Bampilis

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14. What functional groups are present in the compound below?

(Α) Ester and ether (D) Ether and carboxylic acid

(B) Ester and amine (Ε) Ether and ketone

15. Which of the following compounds contains the greatest percentage of oxygen by weight? (Α) C3H6O5Cl b. C3H6O2 (C) C5H10O5 (D) C4H8O3 (Ε) All are equal 16. The first and simplest alkane is: (Α) Ethane (B) Methane (C) C2H2

(D) Methene

(Ε) CCl4

17. Compounds that have the same composition but differ in structural formulas: (Α) Are used for substitution products (B) Are called polymers (C) Are usually alkanes (D) Have the same properties (Ε) Are called isomers 18. Ethene is the first member of the: (Α) Alkane series (B) Alkyne series (C) Saturated hydrocarbons (D) Unsaturated hydrocarbons e. Aromatic hydrocarbons 19. The characteristic group of the organic ester is: (Α) –CO(B) –COOH (C) –CHO (D) –O20. Coke is produced from bituminous coal by: (Α) Cracking (B) Synthesis (C) Substitution

(Ε) –COO-

(D) Destructive distillation

21. An ester can be prepared by the reaction of: (Α) Two alcohols (B) An alcohol and an aldehyde (C) An alcohol and an organic acid (D) An organic acid and an aldehyde (Ε) An acid and a ketone 22. The usual method for preparing carbon dioxide in the laboratory is: (Α) Heating a carbonate (B) Fermentation (C) Reacting an acid and a carbonate (D) Burning carbonaceous materials 23. Slight oxidation of a primary alcohol gives: (Α) a ketone (B) an organic acid (C) an ether

(D) an aldehyde

(Ε) an ester

24. The organic acid that can be made from ethanol is: (Α) acetic acid (B) formic acid (C) C3H7OH (D) Found in bees and ants (Ε) Butanoic acid SAT CHEMISTRY Dr. D. Bampilis

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25. The normal electron configuration for ethyne (acetylene) is: (Α) H:C::C:H (B) H:C:C:H (C) H•C:::C•H (D) H:C:::C:H (Ε) H:C:C:H 26. The atomic structure of the alkane series contains hybrid orbitals designated as: (Α) sp (B) sp2 (C) sp3 (D) sp3d2 (Ε) sp4d3 27. Which of the following statements is the best expression for the sp3 hybridization of carbon electrons? (Α) The new orbitals are one s orbital and three p orbitals (B) The s electron is promoted to the p orbitals (C) The s orbital is deformed into a p orbital (D) Four new and equivalent orbitals are formed (Ε) The s orbital electron loses energy to fall back into a partially filled p orbital 28. The following statements about carbon dioxide are true EXCEPT: (Α) It can be prepared by the action of acid on CaCO3 (B) It is used in fire extinguishers (C) It dissolves in water at room temperature (D) It sublimes rather than melts at 20oC and 1 atm pressure (Ε) It is a product of photosynthesis in plants 29. The structure of the third member of the alkyne series is: (Α) H—C≡C—H (B) H—C≡C—CH3 (C) H—C≡C—CH2CH3 (D) H—C≡C—C≡C—H (Ε) H—C—C—CH=CH2 30. The primary products of hydrocarbon combustion are: (Α) Water and carbon (B) Water and carbon monoxide (C) Water and carbon dioxide (D) Hydrogen and carbon monoxide (Ε) Hydrogen and carbon 31. The production of alkanes from alkenes is accomplished by: (Α) Burning in the presence of water (B) Distillation (C) Methylation (D) Catalytic hydrogenation (Ε) Hydrolysis 32. sp2 hybridization will be found for carbon in: (Α) CH4 (B) C2H4 (C) C2H6

(D) CH3OH

(Ε) CH3OCH3

33. The functional group shown below represents:

(Α) An alcohol (B) An ether (Ε) An organic acid derivative

(D) A ketone

34. Which of the following is the functional group of an ether? SAT CHEMISTRY Dr. D. Bampilis

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(Α) R—OH

(B) R—O—R’

(C)

(D)

(Ε)

35. A triple bond may best be described as: (Α) Two sigma bonds and one pi bond (B) Two sigma bonds and two pi bonds (C) One sigma bond and two pi bonds (D) Three sigma bonds (Ε) Three pi bonds 36. Which substance yields the most energy per gram of sample upon metabolism? (A) carbohydrate (B) fat (C) protein (D) vitamin 37. Which of the following hydrocarbons would be expected to have the highest boiling point? (Α) CH 4 (Β) C3H8 (C) C 4 Η 10 (D) C 5 H 12 (Ε) C6 Η 14 38. When methane, CH4 , burns in excess oxygen, the products would be: (Α) CH4 O 2 (Β) CO+Η 2 Ο (C) CO+CH2 OH (D) CO 2 +Η 2 Ο (Ε) CO 2 +2Η 2 39. The classification of a fat as saturated or unsaturated is based on whether: (A) it can be metabolized by humans. (B) it contains carbon – carbon double bonds. (C) it has twenty or more carbon atoms. (D) it is of animal origin. 40. Which compound is INCORRECTLY matched to the functional group that it contains? (A) CH3 COOH hydroxyl (Β) CH3 OH hydroxyl (C) CH3 CH2 NH2 amine (D) CCI 3 COOH carboxylic acid (Ε) C 6 H5 COOH carboxylic acid 41. Refining of petroleum requires the separation of its components into different fractions. BECAUSE The hydrocarbon chains that make up petroleum have the same basic carbon chain (same number of C's in the parent).

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15. Nuclear Chemistry All of the processes discussed in this section are examples of nuclear reactions, which are different from ordinary chemical reactions. Ordinary chemical reactions involve the exchange and sharing of electrons, while nuclear reactions involve alterations in the very core of an atom; that dense nucleus made up of protons and neutrons.

Radioactivity You will need to be familiar with several types of nuclear reactions and terms related to them to be fully prepared for the SAT II Chemistry test, and in this section we’ll review everything you’ll need to know. The first concept we discuss is radioactivity. Strictly speaking, radioactivity is the spontaneous disintegration of an unstable atomic nucleus and the subsequent emission of radiation. But what makes atoms radioactive to begin with, and what makes them undergo radioactive decay? It turns out that there is a stable ratio of protons to neutrons for each element; for the first 20 elements on the periodic table (hydrogen through calcium), this ratio is 1 proton to 1 neutron, for example. Protons and neutrons in excess of this stable number can be emitted radioactively. Below we have listed examples of the important types of radioactive decay. Alpha decay occurs when the nucleus emits an alpha particle. Alpha particles have a positive charge and are equivalent in size to a helium nucleus, and so they are symbolized as . Alpha particles are the largest radioactive particle emitted. This type of radioactivity results in a decrease in the atomic number by 2 and a decrease in the atomic mass by 4. The equation below shows uranium-234 undergoing alpha decay: Beta decay occurs when the nucleus emits a beta particle. Beta particles have a negative charge and are much smaller than alpha particles. They’re equivalent to high-speed electrons and are symbolized by or . This type of radioactivity causes an increase in the atomic number by 1 but no change in mass number. The equation below represents uranium-233 undergoing beta decay. How does a nucleus, which is composed of only protons and neutrons, eject an electron? A neutron is composed of a proton and an electron fused together. In beta emission, the electron is emitted from the nucleus, while the proton part remains behind, thus increasing the atomic number by 1. Example Complete the balanced equation by determining the missing term. Explanation Remember, the sum of the atomic numbers and the mass numbers must be equal on both sides of the equation. We are looking for a component that has mass number of 80 and an atomic number of 34 (34 protons). Using this information and the periodic table, we can identify the element produced by this beta decay as Se, or selenium. The missing term is Se. And the completed equation is: Gamma decay consists of the emission of pure electromagnetic energy; no particles are emitted during this process, and it is symbolized by equation;

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After beta, positron, or alpha decay, the nucleus is left in a high-energy state, and at this point it will often emit gamma rays, which allows it to relax to its lower-energy ground state. Since gamma rays do not affect charge or mass, they are often not included in nuclear equations. Positron emission occurs when an atom becomes more stable by emitting a positron01e, which is the same size and mass as an electron but has a positive charge. This process converts a proton into a neutron; the positron is emitted and the neutron remains behind in the nucleus, decreasing the atomic number by 1. Often the emission of an alpha or a beta particle creates another radioactive species, which undergoes further radiation/emission in a cascade called a radioactive series. Notice that in the course of all of these types of radioactive decay, neither protons nor neutrons are either created or destroyed: this is due to what’s known as the law of conservation of matter, which states that mass is neither created nor destroyed. So when you see radioactivity equations on the SAT II Chemistry test, one of the most important things to remember is that the sum of the mass numbers and the sum of the atomic numbers must both be equal on both sides of the equation. Example Write the equation for the alpha decay of radium-221.Write the equation for the beta decay of sulfur-35. Explanation The radium-221 atom has atomic number (A) = 88 and mass number (Z) = 221. When an alpha particle is emitted, the atomic number is reduced by 2 and the mass number is reduced by 4. The atomic number of the resulting atom is 86, so the element created as a result of this radioactive decay is radon-217. The sulfur-35 atom has an atomic number of 16 and a mass number of 35. When it undergoes beta decay, the atomic number is increased by 1 and the mass number remains the same. The atomic number of the atom created is 17, so the atom is chlorine-35.

Fission and Fusion There are two main types of nuclear reactions: fusion and fission. In fusion reactions, two light nuclei are combined to form a heavier, more stable nucleus. In fission reactions, a heavy nucleus is split into two nuclei with smaller mass numbers. Both processes involve the exchange of huge amounts of energy: about a million times more energy than that associated with ordinary chemical reactions. In either case, if the new particles contain more stable nuclei, vast quantities of energy are released. Nuclear power plants rely on fission to create vast quantities of energy. For example, U-235 nuclides can be bombarded with neutrons, and the result is lots of energy, three neutrons, and two stable nuclei (Kr-92 and Ba-141). The three neutrons formed can collide with other U-235 atoms, setting off a chain reaction and releasing tons of energy.

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Example Is the following process an example of fission or fusion? Explanation This is an example of fission. Fission occurs when a large nucleus is bombarded by a small particle, such as a neutron. The result is two smaller nuclei and additional neutrons, and a chain reaction process begins.

Half-Lives In discussions of radioactivity, the half-life of an isotope refers to the time it takes for onehalf of the sample to decay. If we start with 100 g of a radioactive substance whose half-life is 15 days, after 15 days 50 g of the substance will remain. After 30 days, 25 g will remain, and after 45 days, 12.5 g remains, and so on. Example A radioactive substance has a half-life of 20 minutes. If we begin with a 500 g sample, how much of the original sample remains after two hours? Explanation The easiest way to attack these questions is to start with the original amount of the sample, then draw arrows representing each half-life. Two hours is 120 minutes, so that’s six halflives. At the end of the stated time period, 7.8 g remains. 500 g 250 g 125 g 62.5 g 31.25 g 15.625 g 7.8125 g

You have to know: Nuclear binding energy and mass defect (mass converted to energy E = mc2) Fusion (joing of nuclei) Fission (spliting of nuclei) 4 2 Alpha particles a or 2 He

( 42 He) 1 C 10

High energy Range: 5 cm in air Stopped by thinner materials (skin, paper) Ionizes gas molecules Beta particles β- or

o 1

e or

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c Low energy Range: 12 m Stopped by 1 cm of Al or a book Ionizes gas molecules

)

Gamma Rays γ High energy radiation

c Stopped by 13 cm Pb Breaks down larger molecules Deflection of radioactive emission (by an electrical field) Other nuclear symbols for subatomic Particles 1 proton 1 p or

1 1

1 neutron 0 n

positron

o 1

e or

Nuclear reactions Control Z – A • a Zdaughter

Zparent

A daughter

A parent

• β Zdaughter

• γ

Zparent 1

A daughter

A parent

Zdaughter

Zparent

A daughter

A parent

Half life t1/2

1 2

• Fraction of original nuclei remaining after n half – lives: ( ) n

1 2

• Fraction of nuclei that has decayed away after n half – lives: 1 ( ) n

n t

n oe

λ: decay constant Radioactive dating 14 6

14 7

o 1

e t1/2

5700

Food preservation Radiotracers – Medicine Radiation therapy

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Exercises 1. The element Lawrencium was first synthesized by the reaction: 25298Cf +115B ___ What products are needed to balance this equation? (A) 42α + 10n (B) 5 0-1e (C) 5 01e (D) 5 10n

258

103Lr

2. Oxygen-15 has a half-life of 9.98 minutes. How much of a 20.0 g sample of oxygen-15 remains after 60.0 minutes? (Α) 0.156 g (Β) 0.312 g (C) 0.625 g (D) 1.25 g (Ε) 2.50 g 3. Complete the balanced equation below. The missing term is: 14 10 6C 4 Be +? 0 (Α) -1 β (B) 0 1 β (C) 4 2 He (D) 0 0 γ (Ε) 1 0 n

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16. Descriptive Chemistry Chemistry of Some Common Substances There will probably be several questions on the SAT II exam that will ask about some common properties of chemicals. The list below constitutes some of the things that everyone should know about chemistry. Group 1A (Alkali Metals) This group consists of the most active metals on the periodic table; these metals react with water at room temperature to form bases. They react readily with acids to produce hydrogen gas and get even more reactive as you move down the family. This makes sense because as you move down the family, there are more energy levels, more shielding, so it’s harder for the nucleus to hold on to the lonely valence electron, and so on. Many drain cleaners contain sodium hydroxide. Group 7A (Halogens) This group contains the most reactive nonmetals on the periodic table, and all of these elements are diatomic. Fluorine is a gas, bromine is a liquid, and iodine is a solid, which makes sense because as the molecules get larger, there are more intermolecular forces to hold them together. Fluorine is the most reactive of the halogens. Chlorine is a very common antibacterial agent, found in bleach and muriatic acid (HCl), and is added to every city’s water supply. Fluorine is the anti-tooth-decay element. Most cities also add fluoride ion to the water supply. Group 8A (Noble Gases) The noble gases are considered the most stable family on the periodic table. Many of these gases appear in signs (such as neon signs). Helium is used to fill balloons because it is much less dense than air. Argon is fairly abundant in our atmosphere. Metals You might recall from our earlier discussions (see “The Structure of Matter”) that metals have a positive center surrounded by a sea of electrons. This sea of electrons makes metallic substances very good conductors of electricity. Alloys are substances that contain a mixture of elements that have metallic properties. An alloy is often much stronger than the individual metal itself. Some of the more common alloys include Brass: mixture of copper and zinc Sterling silver: mixture of silver and copper Steel: mixture of iron and carbon Bronze: mixture of copper, zinc, and other metals Pewter: mixture of tin, copper, bismuth, and antimony Properties of Some Common Gases Hydrogen: H2 is a colorless, odorless gas. It was once used to fill blimps because of its low density, but now helium is used since hydrogen is very flammable. When hydrogen gas is collected in a test tube in the lab, a burning splint inserted into the test tube filled with hydrogen will “bark” as the hydrogen ignites. Oxygen: O2 makes up about 21% of our atmosphere (the other major gases that make up the atmosphere are nitrogen and argon). It is a colorless, odorless gas that is necessary for life and supports combustion reactions. When oxygen is collected in a test tube in the laboratory, a glowing wooden splint will reignite. Carbon dioxide: CO2 is also a colorless, odorless gas that does not support combustion; many fire extinguishers use carbon dioxide to extinguish flames. When carbon dioxide gas is collected in a test tube in the laboratory, a burning wooden splint will go out when placed into the gas. Another common lab test for CO2 is to bubble it into limewater, Ca(OH)2. The clear solution will turn cloudy as calcium carbonate, CaCO3, begins to precipitate. Chlorine: Cl2 is a deadly yellow-green gas. It has often been used as a weapon in warfare.

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Environmental Chemistry Fuels The major sources of energy in the United States are coal, petroleum, and natural gas, all of which are known as fossil fuels. Fossil fuels were formed millions of years ago by the decomposition of animals and plants and thus are in limited supply. We are quickly depleting the available fossil fuels. Coal is solid and is composed of large hydrocarbons and other compounds that contain sulfur, oxygen, and nitrogen. When it’s combusted, the sulfur it contains is converted to SO2, which is an air pollutant. Petroleum is a liquid made up of hundreds of different components, but mostly hydrocarbons. It also contains some compounds that have functional groups containing sulfur, nitrogen, or oxygen. The first step in refining (processing) petroleum is to separate it into fractions based on the different boiling points of its components. Natural gas consists of hydrocarbons in the gas phase, primarily methane (CH4). Air Pollution Air pollution is the contamination of air by a variety of substances, causing health problems and damaging our environment. It has thinned the ozone layer above the earth, exposing us to harmful UV radiation from the sun. Some of the major pollutant gases are listed below. Carbon monoxide: CO is produced from incomplete combustion of all types of natural and synthetic products, including cigarette smoke. When it builds up in high concentrations, it can be very toxic. Cities with heavy traffic problems are known for dangerous CO levels. Carbon dioxide: CO2 is the principal greenhouse gas and is primarily responsible for the greenhouse effect. It can be formed from all types of common human activity, such as burning fuels and even breathing. Chlorofluorocarbons: Chlorofluorocarbons, or CFCs, are used in great quantities in industry, for refrigeration and air-conditioning, and in consumer products. When released into the air, they rise into the stratosphere, where they readily react with the ozone that constitutes the ozone layer, effectively degrading it. Ozone: O3 gas occurs naturally in the upper atmosphere, where it shields the earth from the sun’s dangerous ultraviolet rays. When found at ground level, however, it’s a pollutant. It can cause damage to humans (especially our respiratory system), the environment, and a wide range of natural and artificial materials. Vehicle exhaust and industry waste are major sources of ground-level ozone. Nitrogen oxide and sulfur dioxide: NOx and SOx are major contributors to smog and acid rain. These gases both react with volatile organic compounds to form smog, which can cause respiratory problems in humans. Acid rain can harm vegetation, change the chemistry of river and lake water by lowering the pH so that it’s harmful to animal life, and react with the marble of statues and buildings and decompose them.

Exercises 1. How many of the following gases are characterized by BOTH color AND a distinctive odor? Cl2 CH4 NO2 (A) nοne (B) one (C) two (D) three 2. Which of the following gases is known to shield the earth from harmful ultraviolet radiation? (Α) CO (g) (B) CO 2 (C) CFCs (D) SO 2 (Ε) O3

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3. Α student performed an experiment and a gas was produced. After the gas was collected and tested with a burning splint, a loud popping noise was heard. Which of the following gases was produced? (Α) Hydrogen (Β) Oxygen (C) Carbon dioxide (D) Chlorine (Ε) Methane 4. Alloys are mixtures of metallic substances. Which of the following pairs are matched INCORRECTLY? (Α) Brass - copper and zinc. (Β) Steel - iron and copper. (C) Bronze - copper, zinc and others. (D) Pewter - tin, copper, bismuth, and antimony. (Ε) Sterling silver - silver and copper. 15.5. All of the following statements about carbon dioxide are true EXCEPT: (A) It can be prepared by the action of acid on limestone. (B) It is used to extinguish fires. (C) It dissolves in water at room temperature. (D) It sublimes rather than melts at 20oC and 1 atmosphere pressure. (E) It is less dense than air at a given temperature and pressure. 6. Which statement is INCORRECT? (Α) All of the gases in the atmosphere mix completely unless they react with each other. (Β) The atmosphere of our planet consists of a mixture of gases and various particles in the liquid and solid state. (C) The major gaseous components of our atmosphere are nitrogen, oxygen, and argon. (D) Carbon dioxide, another major component of our atmosphere, has concentrations that are relatively the same everywhere within our atmosphere. (Ε) The amount of moisture in our atmosphere varies with location. 7. The combustion of fuels containing sulfur leads to the production of acid rain. BECAUSE Sulfur oxides form acid solutions in water.

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17. Laboratory Rules for Basic Laboratory Safety 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13.

Safety goggles must be worn at all times in the laboratory. No eating or drinking in the laboratory. Never taste or touch the laboratory chemicals. Always wash your hands before leaving the laboratory. Wear proper clothing—safety glasses, closed-toed shoes, and an apron; tie long hair back and remove all jewelry. Always follow the written directions, and never perform an unauthorized experiment. Always add acid to water. This prevents the acid from spattering. Point heating test tubes away from others and yourself, and heat them slowly. Never return unused chemicals to their original containers. This prevents contamination. Always use a pipette bulb or a pipette to transfer when using a pipette. Never use your mouth. Always use a fume hood when working with toxic substances. Never inhale fumes directly. Never use an open flame near flammable liquids. Dispose of chemicals in the designated disposal site—not in the sink or trash can.

SAT CHEMISTRY Dr. D. Bampilis

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SAT CHEMISTRY Dr. D. Bampilis

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SAT CHEMISTRY Dr. D. Bampilis

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Some Common Lab Techniques Massing solids: When obtaining the mass of solid chemicals, always use some type of weighing paper to protect the pan of the balance. Remember that the mass of the weighing paper must be written down and subtracted from the total weight when you are determining the amount of solid obtained. Measuring liquids: When measuring out a particular volume of a liquid, you must choose an instrument that will measure as accurately as possible. For small quantities it would be appropriate to use a pipette or burette. For larger quantities a graduated cylinder might be appropriate. Remember that beakers are not accurate measuring instruments! Remember always to take measurements of liquids from the bottom of the meniscus. SAT CHEMISTRY Dr. D. Bampilis

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Filtering: When filtering a solid from a mixture by gravity filtration, always weigh the filter paper, fold it, place it in the funnel, and wet it down to hold it in place before beginning the filtering process. After filtering, the solid on the filter paper must be dried and weighed. The initial weight of the filter paper is subtracted to find the mass of the solid obtained. The liquid that comes through the filter paper is known as the filtrate. Color Review One way to identify elements is by performing a simple flame test in the laboratory. When the electrons are heated, they get excited and jump away from the nucleus. As they fall back down, they release energy, often in the form of visible light. Some of the most common colors of flames are listed. You may recognize many of these from fireworks displays! Ion Flame color + 2+ 2+ Li , Sr , Ca Red + Na Yellow + K Purple (pink) 2+ Ba Light green 2+ Cu Blue-green 3+ Fe Gold Many solutions in chemistry also have color, which is often the result of unpaired electrons. Metal ions often are colored. Ion Solution color 2+ Cu Blue 3+ Fe Yellow to orange (rusty) 2+ Ni Green Purple Yellow Orange

Exercises 1. Which technique can be used to determine the number of components in a plant pigment? (A) calorimetry (B) chromatography (C) colorimetry (D) gravimetry 2. A Material Safety Data Sheet (MSDS) provides what type(s) of information about a chemical? I. First aid measures. II. Handling and storage tips. (A) I only (B) II only (C) Both I and II (D) Neither I nor II 3. Compounds of uranium-235 and uranium-238 can be separated from one another by : (A distillation (B) effusion (C) fractional crystallization (D) paper chromatography 4. A student is asked to dispense 24.70 mL of a solution with an uncertainty of less than 0.05 mL. Which item should be used for this task? (A) 50 mL beaker (B) 50 mL buret (C) 50 mL Erlenmeyer flask (D) 50 mL graduated cylinder

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5. Which action should be taken immediately if concentrated sulfuric acid is spilled on the skin? (A) It should be rinsed off with large quantities of running water. (B) It should be neutralized with solid CaCO3. (C) It should be neutralized with concentrated NaOH. (D) The area of the spill should be wrapped tightly with cloth and shown to a health provider. 6. Which instrument would be best suited for use in a volumetric analysis to fin d the unknown molarity of base when titrated with a known acid? (Α) Graduated cylinder (Β) Pipette (C) 250 mL beaker (D) Burette (Ε) Triple beam balance 7. Which of the following laboratory techniques would not be a physical change in components of a mixture? (Α) Chromatography (Β) Precipitation (C) Filtering (D) Distillation (Ε) Evaporation 8. Which of the following mixtures would best be separated by gravity filtration? (Α) Α solid precipitate in a liquid solution. (Β) A mixture of oil and water. (C) Α mixture of solid iron with solid sulfur. (D) Carbon dioxide gas bubbles in a soft drink. (Ε) Α mixture of dyes in a felt-tip pen. 9. Which of the following statements is the most probable explanation for color in solutions? (Α) Solutions with color contain oxygen. (Β) Solutions with color contain metals. (C) Solutions with color are ionically bonded with water. (D) Solutions with color usually contain transition metals with unshared electron pairs. (Ε) Solutions with color have electron configurations that are isoelectronic with the noble gases. 10. Α student mixes 10.0 mL of 0.10 Μ AgNO 3 with excess copper metal. The reaction should produce 0.107 gram of silver; however, the student obtains a mass of 150 grams of silver. Possible explanations for this yield >100% might include: Ι. The student did not subtract the mass of the filter paper before recording results. ΙΙ. The student did not thoroughly dry the sample before massing. Ill. The copper metal did not react completely. (Α) Ι only (B) II only (C) I and II only (D) I and III only (Ε) Ι, ΙΙ, and III

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11. In a neutralization reaction performed in lab, a student mixed 0.20 Μ ΝaΟΗ with 0.10 Μ HCI until the reaction was complete. After the liquid left in the container was dried, which of the following statements must be true? Ι. The student produced a salt and water. All that was left in the container was the salt. ΙI. The total mass of the products in the evaporating dish at the end of the experiment had a lower mass than before heating. ΙΙΙ. The student was left with an ionically bonded, white, crystalline solid. (Α) Ι only (B) II only (C) I and II only (D) I and III only (Ε) Ι, ΙΙ and III all are true

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Practice test 1 For 1 - 4: (Α)CH3CH2CH2OH (D) CH3CH2CH2COOH

(B) CH3CH(Br)CH(Br)CH3 (Ε) CH3CH2CH2CH3

1. CH3CH=CHCH3 + H2

2. 1. + Br2

HBr + 2.

3. CH3CH=CHCH3 + Br2

4. Which compound would most likely turn litmus paper to a red color? For 5 - 8: (Α)Heisenberg Uncertainty principle (B) Pauli Exclusion Principle (C) Schrodinger Wave Equation (D) Hund’s Rule (Ε) Bohr model of the hydrogen atom 5. No two electrons can have the same quantum number because they must have opposite spins. 6. We cannot know the exact location of an electron in space. 7. The electrons will occupy an orbital singly, with parallel spins, before pairing up. 8. The energy changes that an electron may undergo are quantized. For 9 - 12: (Α)H2

(B) CO2

(D) NaCl

(Ε) CH2CH2

9. Contains just one sigma bond 10. Has a bond formed from the transfer of electrons 11. Has an atom that is sp hybridized 12. Is a polar molecule For 13 - 16: (Α)F (B) Li

(D) He

(Ε) Si

13. Shows both the properties of both metals and non-metals 14. Has the greatest ionization energy SAT CHEMISTRY Dr. D. Bampilis

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15. Has the greatest electronegativity 16. Has colored salts that will produce colored aqueous solutions For 17 - 19: (Α)NaC2H3O2

(B) HC2H3O2

(D) NH3

(Ε) HCl

17. Is a salt that will undergo hydrolysis to form a basic solution 18. Will form a coordinate covalent bond with a hydronium ion 19. Is a strong acid For 20 - 22: (Α) q = mcΔT

(B) q = Hvm

(D) D = m/V

(Ε) K = C + 273

20. Can be used to find the mass of an irregularly shaped solid 21. Boyle’s Law 22. Used to find energy gained or lost during a particular phase change. For 23 - 25: (Α)Alpha particle (D) Gamma ray

(B) Beta particle (Ε) Positron

23. Has the greatest mass 24. Has the greatest positive charge 25. Has the same mass and charge as an electron Q Statement I BECAUSE Statement II 26. 12C is an isotope of 14C Because The nuclei of both atoms have the same number of neutrons 27. Ne is an inert gas BECAUSE Ne has a complete octet in its valence shell 28. A solution with a pH of 5 is less acidic than a solution with a pH of 8 BECAUSE A solution with a pH of 5 has 1000 times more hydronium ions than a solution with a pH of 8 SAT CHEMISTRY Dr. D. Bampilis

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29. A reaction with a positive ΔH is considered to be exothermic BECAUSE An exothermic reaction has more heat released than absorbed 30. A voltaic cell spontaneously converts chemical energy into electrical energy BECAUSE A voltaic cell needs an externally applied current to work 31. K is considered to be a metal BECAUSE When K becomes an ion its atomic radius increases 32. At equilibrium the concentration of reactants and products remain constant BECAUSE At equilibrium the rates of the forward and reverse reactions are equal 33. Powdered zinc will react faster with HCl than one larger piece of zinc of the same mass BECAUSE Powdered zinc has less surface area than one larger piece of zinc of the same mass 34. An organic compound with the molecular formula C4H10 can exist as two compounds BECAUSE n-butane and 2-methylpropane are isomers that have the molecular formula of C4H10 35. At STP, 22.4 liters of He will have the same volume as one mole of H2 (assume ideal gases) BECAUSE One mole or 22.4 liters of any gas at STP will have the same mass 36. Halogen molecules can exist as solids, liquids or gases at room temperature BECAUSE As nonpolar molecules are considered by increasing mass the dispersion forces between them increases 37. Hydrocarbons will dissolve in water BECAUSE Substances that have the same polarity are miscible and can dissolve each other 38. Ammonia has a trigonal pyramidal molecular geometry BECAUSE Ammonia has a tetrahedral electron pair geometry with the three atoms bonded to the central atom 39. AlCl3 is called aluminum trichloride SAT CHEMISTRY Dr. D. Bampilis

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BECAUSE Prefixes are used when naming covalent compounds 40. When a Li atom reacts and becomes an ion, the Li atom can be considered to be a reducing agent BECAUSE The Li atom lost an electron and was oxidized 41. 117 grams of NaCl are dissolved in water to make 500 mL of solution. Water is then added to this solution to make a total of one liter of solution. The final molarity of the solution will be (Α)4 M (B) 2 M (C) 1 M (D) 0.5 M (Ε) 0.585 M 42. How many pi bonds are in the molecule 2-butyne, CH3C CCH3? (Α)1 (B) 2 (C) 4 (D) 6 (Ε) 10 43. How many atoms lie in a straight line in the molecule 2-butyne, CH3C CCH3? (Α)10 (B) 8 (C) 6 (D) 4 (Ε) 2 44. A solution of a weak acid, HA, has a concentration of 0.100 M. What are the concentration of hydronium ion and the pH of this solution if the Ka value for this acid is 1.0 x 10-5? (Α)1.0 x 10-3 and pH = 11 (B) 1.0 x 10-6 and pH = 6 (C) 1.0 x 10-4 and pH = 8 (D) 3.0 x 10-4 and pH = 4 (Ε) 1.0 x 10-3 and pH = 3 45. Given the reaction at STP: Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g), how many liters of H2(g) can be produced from the reaction of 12.15 g Mg and excess HCl(aq)? (Α)2.0 L (B) 4.0 L (C) 11.2 L (D) 22.4 L (Ε) 44.8 L 46. A student performed a single titration using 2.00 M HCl to completely titrate 40.00 mL of 1.00 M NaOH. If the initial reading on the buret containing HCl was 2.05 mL, what will be the final reading? (Α)82.05 mL (B) 42.05 mL (C) 20.00 mL (D) 10.00 mL (Ε) 22.05 mL 47. Which of the following was NOT a conclusion of Rutherford’s gold foil experiment? (Α)The atom is mainly empty space (B) The nucleus has a negative charge (C) The atom has a dense nucleus (D) Alpha particles can pass through a thin sheet of gold foil (Ε) All of the above are correct regarding the gold foil experiment 48. In a reaction the potential energy of the reactants is 40 kJ/mol, the potential energy of the products is 10 kJ/mol and the potential energy of the activated complex is 55 kJ/mol. What is the activation energy for the reverse reaction? (Α)45 kJ/mol (B) -30 kJ/mol (C) 15 kJ/mol (D) 35 kJ/mol (Ε) -55 kJ/mol SAT CHEMISTRY Dr. D. Bampilis

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49. Which reactions would form at least one solid precipitate as a product? Assume aqueous reactants. i. AgNO3 + NaCl NaNO3 + AgCl ii. Pb(NO3)2 + 2KI PbI2 + 2KNO3 iii. 2NaOH + H2SO4 Na2SO4 + 2H2O (Α)i only (B) ii only (C) iii only (D) i and ii only (Ε) ii and iii only 50. What is the mass action equation for: 2A(aq) + B(aq) ⇌ 3C(aq) + D(s)? (Α)Κ =

(B) Κ =

(D) Κ =

(Ε) Κ =

51. If the equilibrium constant for a reverse reaction is 9.0 x 10-4, what is the equilibrium constant for the forward reaction? (Α)3.0 x 10-2 (B) -3.0 x 10-2 (C) -9.0 x 10-2 (D) 1 / 9.0 x 10-4 (Ε) 1 / -9.0 x 10-4 52. A compound’s composition by mass is 50% S and 50% O. What is the empirical formula of this compound? (Α)SO (B) SO2 (C) S2O (D) S2O3 (Ε) S3O4 53. What percentage of the total mass of KHCO3 is made up by nonmetallic elements? (Α)17% (B) 83% (C) 61% (D) 20% (Ε) 50% 54. Which aqueous solution is expected to have the highest boiling point? (Α)0.2 m CaCl2 (B) 0.2 m NaCl (C) 0.1 m AlCl3 (D) 0.2 m CH3OH (Ε) 0.2 m NaC2H3O2 55. Which of the following solids are known to undergo sublimation? i. CO2 ii. I2 iii. Napthalene (Α)i only (B) ii only (C) i and ii only (D) ii and iii only iii

(Ε) i, ii and

56. Which of the following demonstrates a decrease in entropy? (Α)Dissolving a solid into a solution (B) An expanding universe (C) Burning a log in a fireplace (D) Raking up leaves into a trash bag (Ε) Spilling a glass of water 57. Which of the following substances is/are liquid(s) at room temperature? i. Hg ii. Br2 iii. Si (Α)i only (B) ii only (C) i and ii only (D) ii and iii only (Ε) i, ii and iii 58. Which of the following would be considered to be unsafe in a laboratory setting? (Α)Using a test tube holder to handle a hot test tube (B) Tying one’s long hair back before experimenting (C) Wearing open-toed shoes SAT CHEMISTRY Dr. D. Bampilis

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(D) Pouring liquids while holding the reagent bottles over the sink (Ε) Working under a fume hood 59. A sample of a gas at STP contains 3.01 x 1023 molecules and has a mass of 20.0 grams. This gas would: (Α)have a molar mass of 20.0 g/mol and occupy 11.2 liters (B) occupy 22.4 liters and have a molar mass of 30.0 g/mol (C) occupy 22.4 liters and have a molar mass of 20.0 g/mol (D) have a molar mass of 40.0 g/mol and occupy 33.6 liters (Ε) have a molar mass of 40.0 g/mol and occupy 11.2 liters 60. Given the reaction: Ca(s) + Cl2(g) CaCl2(s), when 80 g Ca (molar mass is 40) is reacted with 213 g Cl2 (molar mass is 71) one will have: (Α)40 g Ca excess (B) 71 g Cl2 excess c. 293 g CaCl2 formed (D) 133 g CaCl2 formed (Ε) 113 g CaCl2 formed 61. A student performed an experiment to determine the solubility of a salt at various temperatures. The data from the experiment can be seen below: Trial Temp (oC) Solubility in 100 g water 1 20 44 2 30 58 3 40 67 4 50 62 5 60 84 Which trial seems to be in error? (Α)1 (B) 2 (C) 3 (D) 4 (Ε) 5 62. Given the following reaction at equilibrium: 3H2(g) + N2(g) ⇌ 2NH3(g) + heat energy, which of the following conditions would shift the equilibrium of this reaction so that the formation of ammonia is favored? (Α)Increasing the pressure of the reaction (B) Heating the reaction (C) Removing hydrogen gas from the reaction (D) Adding more ammonia to the reaction (Ε) Removing nitrogen gas from the reaction 63. Given equal conditions, which gas below is expected to have the greatest density? (Α)H2 (B) Ne (C) Ar (D) H2S (Ε) Cl2 64. Given equal conditions, which gas below is expected to have the greatest rate of effusion? (Α)H2 (B) Ar (C) Kr (D) F2 (Ε) Cl2 65. Ideal gases: (Α)Have forces of attraction between them (C) Never travel with a straight line motion (Ε) Have low masses and are spread far apart SAT CHEMISTRY Dr. D. Bampilis

(B) Are always linear in shape (D) Have molecules that are close together

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66. Which substance will combine with oxygen gas to produce a greenhouse gas? (Α)Na (B) S (C) H2 (D) Ne (Ε) C 67. Which general formula below represents that of an organic ester? (Α)R—OH (B) R—COOH (C) R—O—R (D) R—COO—R

(Ε) R—CO—R

68. When an alkaline earth metal, M, reacts with oxygen the formula of the compound produced will be: (Α)M2O (B) MO (C) M2O3 (D) MO2 (Ε) M3O4 69. A catalyst can change the: (Α)Heat of reaction and the potential energy of the reactants (B) Heat of reaction and the time it takes the reaction to proceed (C) Activation energy of the reverse reaction and the potential energy of the activated complex (D) Potential energy of the reactants and the time it takes the reaction to proceed (Ε) Activation energy of the forward reaction and the potential energy of the products 70. A neutral atom has a total of 17 electrons. The electron configuration in the outermost principle energy level will look closest to: (Α)1s22s22p5 (B) 3s53p2 (C) s2p5 (D) s2p8d7 (Ε) sp7 71. Given a 22.4 liter sample of helium gas at STP, if the temperature is increased by 15 degrees Celsius and the pressure changed to 600 torr, what would the new volume be? (Α) (760)(22.4)(15) / (273)(600) (B) (273)(600)(288) / (760)(22.4) (C) (760)(22.4)(15) / (600) (D) (760)(22.4)(288) / (273)(600) (Ε) (273)(600) / (760)(22.4)(288) 72. Which of the following are correct about the subatomic particles found in 37Cl1-? i. 21 neutrons ii. 17 protons iii. 16 electrons (Α)ii only (B) iii only (C) i and ii only (D) i and iii only (Ε) ii and iii only 73. A hydrated blue copper (II) sulfate salt with a formula YCuSO4•XH2O is heated until it is completely white in color. The student who performed the dehydration of this salt took note of the mass of the sample before and after heating and recorded it as follows: Mass of hydrated salt = 500 g Mass of dehydrated salt = 320 g What is the value of “X” in the formula of the hydrated salt? (Α)1 (B) 2 (C) 4 d. 5 (Ε) 10 74. Which of the following oxides can dissolve in water to form a solution that would turn litmus indicator red in color? (Α)MgO (B) K2O (C) CO2 (D) ZnO (Ε) H2O SAT CHEMISTRY Dr. D. Bampilis

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75. The process in which water vapor changes phase to become liquid is called: (Α)Deposition (B) Sublimation (C) Vaporization (D) Fusion (Ε) Condensation 76. What is the value for ΔH for the reaction N2O4 2NO2 given the following: 2NO2 N2 + 2O2 ΔH = -16.2 kcal N2 + 2O2 N2O4 ΔH = +2.31 kcal (Α)+13.89 kcal (B) +18.51 kcal (C) +37.42 kcal (D) -13.89 kcal (Ε) -18.51 kcal 77. A liquid will boil when: (Α)Enough salt has been added to it (B) The vapor pressure of the liquid is equal to the atmospheric or surrounding pressure (C) The vapor pressure of the liquid reaches 760 mmHg (D) Conditions favor the liquid’s molecules to be closer together (Ε) It has been brought up to a higher elevation 78. A conductivity experiment is set up with a light bulb and five beakers of 0.1 M solutions of the substances below. Which solution would allow the bulb to glow the brightest? (Α)C6H12O6 (B) HCl (C) SiO2 (D) HC2H3O2 (Ε) CH3OH 79. Which of the following represents a correctly balanced half-reaction? (Α)Cl2 + 2e- Cl(B) 2e- + Fe Fe2+ (C) O2 2e- + 2O2(D) Al3+ Al + 3e(Ε) 2H+ + 2e- H2 80. A student prepares for an experiment involving a voltaic cell. Which of the following is needed the least to perform the experiment? (Α)Buret (B) Salt bridge (C) Strip of zinc metal (D) Copper wire (Ε) Solution of zinc sulfate 81. When the equation: __C3H8 + __O2 __CO2 + __H2O is balanced using the lowest whole number coefficients, the coefficient before O2 will be: (Α)1 (B) 2.5 (C) 5 (D) 10 (Ε) 13 82. Which nuclear equation below demonstrates beta decay? (Α) 238U 234Th + X (B) 1H + X 3H (C) 14N + X 17O + 1H (D) 234Pa 234U + X (Ε) None of the above demonstrates beta decay 83. Which of these processes could be associated with the following reaction: 2H2O O2? i. electrolysis ii. neutralization iii. decomposition (Α)i only (B) iii only (C) i and iii only (D) i and ii only (Ε) ii and iii only SAT CHEMISTRY Dr. D. Bampilis

2H2 +

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84. The following reaction occurs in a beaker: Ag+(aq) + Cl-(aq) sodium chloride were added to this beaker: (Α)The solubility of the sodium chloride would decrease (B) The reaction would shift to the left (C) The concentration of silver ions in solution would increase (D) The solubility of the silver chloride would decrease (Ε) The equilibrium would not shift at all

AgCl(s). If a solution of

85. How many atoms are represented in the equilibrium Pb(NO3)2 + 2KI 2KNO3? (Α)5 (B) 12 (C) 13 (D) 18 (Ε) 26

SAT CHEMISTRY Dr. D. Bampilis

PbI2 +

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Practice test 2 For 1 - 4: (Α)Anions

(B) Cations

(D) Isotope

(Ε) Atom

1. a positive ion 2. an atom of the same element that differs by the number of neutrons 3. cannot be broken down chemically 4. will migrate through the salt bridge to the anode half cell For 5 - 7: (Α)calorimeter (Ε) Bunsen burner

(B) Geiger counter

(D) Funnel

5. used to detect radioactivity 6. used to deliver acids and bases in a titration 7. can be lined with moist filter paper to catch insoluble solids For 8 - 10: (Α)Arrhenius acid (D) Lewis base

(B) Arrhenius base (C) Lewis acid (Ε) Bronsted-Lowry acid

8. yields hydroxide ions as the only negative ions in solution 9. electron pair acceptor 10. proton donor For 11 - 14: (Α)purple solution (D) silver-gray liquid

(B) brown-orange liquid (C) green gas (Ε) yellow-orange when burned in a flame

11. potassium permanganate 12. sodium salt 13. chlorine 14. mercury SAT CHEMISTRY Dr. D. Bampilis

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For 15 - 17: (Α)Eo is positive (D) Keq is greater than 1

(B) ΔS is negative (Ε) Ka is very large

15. indicates a strong acid 16. a reaction is nonspontaneous 17. less chaos, disorder, and randomness For 18 - 21: (Α)alkali metals (D) halogens

(B) alkaline earth metals (Ε) noble or inert gases

18. group 1 19. group 10 20. contains elements in the solid, liquid, and gas phase 21. will form chlorides with the formula MCl2 For 22 - 25: (Α)1+ (B) 1-

(D) 2+ (Ε) 3+

22. oxidation number of O in H2O2 23. oxidation number of F in HF 24. oxidation number of O in O3 25. oxidation number of calcium in calcium phosphate Q Statement I BECAUSE Statement II 26. Methane is defined as a compound BECAUSE Methane can be broken down chemically 27. The burning of a piece of paper is a physical change BECAUSE Once burned, the chemical properties of the paper remain the same 28. -273 degrees Celsius is also known as absolute zero BECAUSE C = K + 273 SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 115

29. The relationship between pressure and volume is considered to be an inverse relationship BECAUSE As pressure increases on a gas, the volume of the gas will decrease 30. A liquid can boil at different temperatures BECAUSE The atmospheric (or surrounding) pressure can vary 31. Bromine has an atomic mass of 79.9 BECAUSE About 50% of all bromine atoms are 79Br and the other 50% are 81Br 32. Excited tungsten atoms will give off light energy BECAUSE As the excited electrons return to their ground state, they emit energy in the form of light 33. As you go from left to right across the Periodic Table, the elements tend to become more metallic in character BECAUSE As you go from left to right across the Periodic Table the elements tend to lose electrons 34. The bonds found in a molecule of N2 are nonpolar covalent BECAUSE There is an equal sharing of electrons between the nitrogen atoms 35. The empirical formula of C6H12O6 is CH2O BECAUSE The empirical formula shows the lowest ratio of the elements present in the molecular formula 36. A solution of NaCl will conduct electricity BECAUSE NaCl will not form ions in solution 37. Increasing the concentration of reactants will cause a reaction to proceed faster BECAUSE More reactants lower the activation energy of a reaction 38. Cl- is the conjugate base of HCl BECAUSE A conjugate base is formed once a Bronsted - Lowry acid accepts a proton 39. F2

2F- + 2e- is a correctly written half reaction

SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 116

BECAUSE This half reaction must demonstrate proper conservation of mass and charge 40. Ethane is considered to be a saturated hydrocarbon BECAUSE Ethene has a triple bond 41. Which of the following would not be attracted or deflected while traveling through an electric field? i. Gamma ray ii. Beta particle iii. neutron (Α)i only (B) ii only (C) i and ii only (D) i and iii only (Ε) i, ii, and iii 42. Which substance below is resonance stabilized by delocalized pi electrons? (Α)Benzene (B) Hydrochloric acid (C) Hydrogen gas (D) Methane (Ε) Potassium bromide 43. Which of the following is true about a solution that has [OH-] = 1.0 x 10-6 M? (Α)The pH is 8 and the solution is acidic (B) The [H+] = 1.0 x 10-8 M and the solution is basic (C) The pH is 6 and the solution is acidic (D) The [H+] = 1.0 x 10-6 M and the solution is basic (Ε) The [H+] = 1.0 x 10-14 M and the solution is neutral 44. What will be the products of the following double replacement reaction? (NH4)3PO4 + Ba(NO3)2 (Α)Ammonium nitrate and barium nitrate (B) Barium nitrate and ammonium phosphate (C) Barium phosphate and sodium nitrate (D) Ammonium nitrate and barium phosphate (Ε) Ammonium nitrate and barium nitrate 45. Which Ka value is that of an acid that is the weakest electrolyte? (Α)1.7 x 10-7 (B) 2.7 x 10-8 (C) 6.6 x 10-10 (D) 4.9 x 10-3 (Ε) 5.2 x 10-4 46. A student performs a titration using 1.00 M NaOH to find the unknown molarity of a solution of HCl. The student records the data as shown below. What is the molarity of the solution of HCl? Base: final buret reading 21.05 mL Base: initial buret reading 6.05 mL mL of base used Acid: final buret reading 44.15 mL Acid: initial buret reading 14.15 mL mL of acid used (Α)0.75 M (B) 0.50 M (C) 0.25 M (D) 0.10 M (Ε) 2.00 M SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 117

47. Which of the following is not a synthetic polymer? (Α)Polyvinyl chloride (B) Plastic (C) Polystyrene (Ε) Cellulose

(D) Polyethylene

48. Which process is represented from area 1 to area 3 on the following phase diagram?

(Α)Evaporation (Ε) Sublimation

(B) Deposition

49. What is the molar mass of Ca3(PO4)2? (Α)310 g/mol (B) 154 g/mol (C) 67 g/mol

(D) 83 g/mol

(D) Freezing

(Ε) 115 g/mol

50. What is the percent composition of oxygen in C6H12O6 (molar mass = 180)? (Α)25% (B) 33% (C) 40% (D) 53% (Ε) 75% 51. The following reaction occurs at STP: 2H2O(l) 2H2(g) + O2(g). How many liters of hydrogen gas can be produced by the breakdown of 72 grams of water? (Α)5.6 liters (B) 11.2 liters (C) 22.4 liters (D) 44.8 liters (Ε) 89.6 liters 52. What is the mass-action expression for the following reaction at equilibrium? 2W(aq) + X(l) ⇌ 3Y(aq) + 2Z(s) (Α)

(B)

(C)

(D)

(Ε)

53. Which statement best describes the bonding found in formaldehyde, CH2O? (Α)The carbon atom is sp hybridized (B) There are three sigma bonds and one pi bond present (C) The bonding gives the molecule a tetrahedral shape (D) The bonds between the atoms are ionic bonds (Ε) All of the bonds are nonpolar bonds 54. What is the molarity of a solution that has 29.25 grams of NaCl dissolved to make 1.5 L of a solution? (Α)19.5 M (B) 3.0 M (C) 1.75 M (D) 0.33 M (Ε) 1.0 M SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 118

55. A sample of a gas at STP contains 3.01 x 1023 molecules and has a mass of 22.0 grams. This gas is most likely: (Α)CO2 (B) O2 (C) N2 (D) CO (Ε) NO 56. What is the value of ΔH for the reaction X + 2Y 2Z? W + X 2Y ΔH = -200 kcal 2W + 3X 2Z + 2Y ΔH = -150 kcal (Α)-550 kcal (B) +50 kcal (C) -50 kcal (D) -350 kcal

(Ε) +250 kcal

57. Which fraction would be used to find the new volume of a gas at 760 torr under its new pressure at 900 torr if the temperature is kept constant? (Α)900 / 760 (B) 1.18 (C) 760 / 900 (D) 658.7 / 798.7 (Ε) 798.7 / 658.7 58. Which of the following aqueous reactions forms a salt that will precipitate out of solution? (Α)HCl + NaOH (B) KBr + NaCl (C) AgNO3 + MgCl2 (D) CaCl2 + KI (Ε) NaNO3 + HC2H3O2 59. Which system at equilibrium will not be influenced by a change in pressure? (Α)3O2(g) ⇌ 2O3(g) (B) N2(g) + 3H2(g) ⇌ 2NH3(g) (C) 2NO2(g) ⇌ N2O4(g) (D) H2(g) + I2(g) ⇌ 2HI(g) (Ε) 2W(g) + X(g) ⇌ 3Y(g) + 2Z(g) 60. The organic reaction: C2H6 + Cl2 HCl + C2H5Cl is best described as: (Α)a substitution reaction (B) an addition reaction (C) an esterification (D) a dehydration synthesis (Ε) a fermentation 61. Enough CaSO4(s) is dissolved in water at 298 K to produce a saturated solution. The concentration of Ca2+ ions is found to be 3.0 x 10-3 M. The Ksp value for CaSO4 will be (Α)6.0 x 10-6 (B) 9.0 x 10-6 (C) 6.0 x 10-3 (D) 9.0 x 10-3 (Ε) 3.0 x 10-3 62. Which of the following statements is not true about acid rain? (Α)Acid rain will erode marble statues (B) Acid rain can change the pH of lakes and streams (C) Acid rain can be formed from carbon dioxide (D) Acid rain creates holes in the ozone layer (Ε) Acid rain can be formed from the gases SO2 and SO3 63. Which mole sample of the solids below is best for melting a 500-gram sheet of ice on a sidewalk? (Α)NaCl (B) CaCl2 (C) KBr (D) AgNO3 (Ε) NaC2H3O2 64. Given this reaction that occurs in plants: 6CO2 + 6H2O C6H12O6 + 6O2, if 54 grams of water are consumed by the plant, how many grams C6H12O6 (molar mass = 180) can be made? Assume an unlimited supply of CO2. SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 119

(Α)54 g

(B) 180 g

(D) 3 g

(Ε) 90 g

65. Which of the following will not be changed by the addition of a catalyst to a reaction at equilibrium? i. The point of equilibrium ii. The heat of reaction, ΔH iii. The potential energy of the products (Α)i only (B) ii only (C) i and ii only (D) ii and iii only (Ε) i, ii and iii 66. According to the reaction Pb(s) + S(s) PbS(s), when 20.7 grams of lead are reacted with 6.4 grams of sulfur (Α)There will be an excess of 20.7 grams of lead (B) The sulfur will be in excess by 3.2 grams (C) The lead and sulfur will react completely without any excess reactants (D) The sulfur will be the limiting factor in the reaction (Ε) There will be an excess of 10.35 grams of lead 67. Which of the following statements is not part of the kinetic molecular theory? (Α)The average kinetic energy of gas molecules is proportional to temperature (B) Attractive and repulsive forces are present between gas molecules (C) Collisions between gas molecules are perfectly elastic (D) Gas molecules travel in a continuous, random motion (Ε) The volume that gas molecules occupy is minimal compared to the volume in which the gas is contained 68. A student is performing an experiment where a blue salt is being heated to dryness in order to determine the percent of water in the salt. Which pieces of laboratory equipment would be used to help determine this percentage? i. A crucible and cover ii. Tongs iii. A triple beam balance (Α)ii only (B) iii only (C) i and iii only (D) d. ii and iii only (Ε) i, ii and iii 69. Which of the following is considered to be a dangerous procedure in the laboratory setting? (Α)Pouring all liquids, especially acids and bases, over the sink (B) Wearing goggles (C) Pushing glass tubing, thermometers, or glass thistle tubes through a rubber cork (D) Pointing the mouth of a test tube that is being heated away from you and others (Ε) Knowing where the fire extinguisher and eyewash stations are located 70. Given a 4-gram sample of each H2(g) and He(g), each in separate containers, which of the following statements is true? (Assume STP) (Α)The sample of hydrogen gas will occupy 44.8 liters and the sample of helium will contain 6.02 x 1023 molecules (B) The sample of hydrogen gas will occupy 22.4 liters and the sample of helium will contain SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 120

3.02 x 1023 molecules (C) The sample of hydrogen gas will occupy 44.8 liters and the sample of helium will contain 1.202 x 1024 molecules (D) The sample of helium will occupy 44.8 liters and the sample of hydrogen gas will contain 6.02 x 1023 molecules (Ε) None of the above statements is correct 71. The diagram shows a solid being heated from below its freezing point. Which line segment shows the gas and the liquid phases existing at the same time?

(Α)A

(B) B

(D) D

(Ε) E

72. Which of the following statements is/are correct regarding molecular geometries? i. CH4 is trigonal pyramidal ii. BF3 is trigonal planar iii. XeF6 is tetrahedral (Α)i only (B) ii only (C) iii only (D) i and iii only (Ε) i, ii and iii 73. When Uranium-238 undergoes alpha decay and then one beta decay, the resulting isotope is: (Α)Th-234 (B) U-234 (C) Pa-234 (D) Th-230 (Ε) Ra-226 74. Which compound is matched up with its correct name? (Α)CO—monocarbon monoxide (B) CaF2—calcium difluoride (C) CCl4—carbon tetrachloride (D) PCl3—potassium trichloride (Ε) TiF4—tin(IV) fluoride 75. Of the statements below, which best explains why CH4 is a gas at STP, while C8H18 is a liquid and C20H42 is a solid? (Α)C20H42 has the greatest ionic interaction between its molecules (B) C20H42 has a greater amount of hydrogen bonding than CH4 or C8H18 (C) There is a more dipole-dipole interaction between molecules of greater mass (D) CH4 has the greatest intermolecular forces while C20H42 has the least (Ε) There are more Van der Waals (dispersion) forces between nonpolar molecules that are greater in mass 76. Which of the gases listed below would not be collected via water displacement? (Α)CO2 (B) CH4 (C) O2 (D) NH3 (Ε) H2 77. Which scientist and discovery are not correctly paired? (Α)Millikan / neutron (B) Rutherford / nucleus (C) Charles / relationship between temperature and volume SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 121

(D) Curie / radioactivity

(Ε) Mendeleyev / periodic table

78. Which of the following situations demonstrate(s) an increase in entropy? i. Dissolving a salt into water ii. Sublimation iii. Heating up a liquid (Α)i only (B) i and ii only (C) ii and iii only (D) i and iii only (Ε) i, ii, and iii 79. How many moles of a gas are present in a closed empty soda bottle that has a volume of 2.0 L at 22oC and a pressure of 1.05 atm? (Α)

(B)

(D)

(Ε)

(C)

80. Which set of conditions below guarantees that a reaction will be spontaneous? (Α)ΔH(+) and ΔS(-) (B) ΔH(-) and ΔS(+) (C) ΔH(+) and ΔS(+) at low temp (D) ΔH(-) and ΔS(-) at high temp (Ε) ΔG(+) 81. How many moles of electrons are transferred in the following reaction? Ce3+ + Pb Ce + Pb4+ (Α)14 (B) 12 (C) 7 (D) 24 (Ε) 3 82. A gas is confined in the manometer as shown below.

The stopcock is then opened and the highest level of mercury inside the tube moved to a level that is 80 mm above its lowest level. What is the pressure of the gas? The atmospheric pressure is 760 mm Hg. (Α)80 mmHg (B) 160 mmHg (C) 680 mmHg (D) 840 mmHg (Ε) The pressure cannot be determined 83. Which statement best describes the density and rate of effusion of the following gases? NO2 C2H6 Kr Xe F2 (Α)Fluorine has the lowest density and the lowest rate of effusion (B) Xenon has the greatest rate of effusion and the lowest density (C) Krypton has the lowest density and the greatest rate of effusion SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 122

(D) Ethane has the greatest rate of effusion and the lowest density (Ε) Nitrogen dioxide has the highest density and the greatest rate of effusion 84. Which indicator is correctly paired up with its proper color if it were added to a base? i. Litmus - blue ii. Phenolphthalein - pink iii. Methyl orange - yellow (Α)i only (B) ii only (C) iii only (D) i and iii only (Ε) i, ii and iii 85. Which structure below demonstrates a violation of the octet rule? (Α)

(B)

(C)

(D)

(Ε)

SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 123

Practice test 3 For 1 - 4:

(Α)Sublimation (Ε) Melting 1. Phase 1 Phase 2 2. Phase 3

Phase 1

3. Phase 1

Phase 3

4. Phase 2

Phase 3

For 5 - 8: (Α)R-COOH

(B) Deposition

(B) R-CHO

(D) R-COO-R

(D) Condensatio

(Ε) R-CO-NH2

5. can be neutralized with a base 6. could be named 2-pentanone 7. amide functional group 8. aldehyde functional group For 9 - 11: (Α)6.02 x 1023 molecules (D) 2.0 moles

(B) 11.2 liters (Ε) 5 atoms

9. 88 grams of CO2(g) at STP 10. 1 molecule of CH4 11. 32 grams of SO2 gas at STP For 12 - 15: (Α) Alkali metal

(B) Alkaline earth metal

SAT CHEMISTRY Dr. D. Bampilis

(D) Halogen

(Ε) Noble gas

12. reacts most vigorously with water 13. is chemically inert 14. has the highest first ionization energy in its period 15. forms ions with a 2+ charge For 16 - 18: (Α)blue

(B) red

(D) colorless

(Ε) orange

16. phenolphthalein in base 17. litmus in acid 18. phenolphthalein in acid For 19 - 22: (Α)Boyle’s law (B) Charles’s law (C) Ideal gas equation (D) Combined gas law (Ε) Dalton’s law of partial pressures 19. Ptotal = P1 + P2 + P3 + … 20. P1V1 = P2V2 21. PV = nRT 22. P1V1/T1 = P2V2/T2 For 23 - 25: (Α)

(B)

(D)

(Ε)

23. has a charge of 2+ 24. has the lowest mass 25. has the greatest mass Q Statement I BECAUSE Statement II 26. The double and single bonds in benzene are subject to resonance BECAUSE Benzene has delocalized pi electrons that stabilize its structure

SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 125

27. The element with an electron configuration of [He] 2s1 has a larger atomic radius than fluorine BECAUSE The element with an electron configuration of [He] 2s1 has a greater nuclear charge than fluorine 28. 1 m NaCl(aq) will have a higher boiling point than that of 1 m CaCl2(aq) BECAUSE 1 mole of NaCl yields 3 moles of ions in solution 29. Neutrons and protons are classified as nucleons BECAUSE Neutrons and protons are both located in the principal energy levels of the atom 30. HCl is considered to be an acid BECAUSE HCl is a proton donor 31. Powdered zinc reacts faster with acid than a larger piece of zinc BECAUSE Powdered zinc has a greater surface area 32. NH3 can best be collected by water displacement BECAUSE NH3 is a polar substance 33. At 1 atm, pure water can boil at a temperature less than 273 K BECAUSE Water boils when the vapor pressure of the water is equal to the atmospheric pressure 34. An exothermic reaction has a negative value for ΔH BECAUSE In an exothermic reaction the products have less potential energy than the reactants 35. As pressure on a gas increases, the volume of the gas decreases BECAUSE Pressure and volume have a direct relationship 36. The addition of H2 to ethene will form an unsaturated compound called ethane BECAUSE Ethane, has as many hydrogen atoms bonded to the carbon atoms as possible 37. AgCl is insoluble in water BECAUSE All chlorides are soluble in water except for those of silver, lead and mercury SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 126

38. ΔS will be positive in value as vaporization occurs BECAUSE Vaporization increases the order of the molecules entering the gas phase 39. Pure water has a pH of 7 BECAUSE The number of H+ ions is equal to the number OH- ions 40. CH3CH2—OH and CH3—O—CH3 are isomers BECAUSE CH3CH2—OH and CH3—O—CH3 have the same molecular formula but different structures 41. One mole of each of the following substances is dissolved in 1.0 kg of water. Which solution will have the lowest freezing point? (Α)NaC2H3O2 (B) NaCl (C) MgCl2 (D) CH3OH (Ε) C6H12O6 42. Which of the following equations is/are properly balanced? i. Cl2 + 2NaBr Br2 + 2NaCl ii. 2Na + O2 Na2O iii. 2K + 2H2O (Α)i only (B) ii only (C) iii only (D) i and iii only iii

H2 + 2KOH (Ε) i, ii and

43. Propane and oxygen react according to the equation: C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(g). How many grams of water can be produced from the complete combustion of 2.0 moles of C3H8(g)? (Α)144.0 (B) 82.0 (C) 8.0 (D) 44.8 (Ε) 22.4 44. A compound was analyzed and found to be composed of 75% carbon and 25% hydrogen. What is the empirical formula of this compound? (Α)C2H4 (B) CH4 (C) CH3 (D) CH2 (Ε) CH 45. Which compound below has a bent molecular geometry? (Α)H2SO4 (B) CH4 (C) CO2 (D) H2S

(Ε) C2H2

46. Of the equipment listed below, which one would require you to read a meniscus? (Α)100 mL beaker (B) 500 mL flask (C) Watch glass (D) 50 mL buret (Ε) Trough 47. Given the following reaction at equilibrium: Fe3+(aq) + SCN-(aq) ⇌ FeSCN2+(aq). Which of these would shift the equilibrium to the left? (Α)Adding FeCl3 to the reaction (B) Adding NH4SCN to the reaction (C) Increasing the pressure on the reaction (D) Adding a catalyst 2+ (Ε) Adding FeSCN (aq) to the reaction

SAT CHEMISTRY Dr. D. Bampilis

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48. Each of the elements listed below is placed in water. Which one will react violently with the water? (Α)Na (B). Fe (C) Cu (D) Au (Ε) Ne 49. Which letter in the boxes below has a value of 7? isotope p n e16 O A 13 C B 23 Na C 10 B (Α)A

(B) B

(D) D

Mass #

Atomic # E

D (Ε) E

50. Which unit is paired incorrectly? (Α)Torr and pressure (B) Mass and grams (C) Heat energy and kilopascals (D) Volume and milliliter (Ε) Temperature and Kelvin 51. Which amount of Pb(NO3)2, when added to enough water to make 1 liter of solution, will produce a solution with a molarity of 1.0 M? (Α) 144 grams (B) 331 grams (C) 317 grams (D) 0.003 moles (Ε) 0.5 moles 52. Enough AgCl(s) is dissolved in water at 298 K to produce a saturated solution. The concentration of Ag+ ions found to be 1.3 x 10-5 M. The Ksp value for AgCl will be: (Α) 2.6 x 10-10 (B) 1.3 x 10-10 (C) 1.3 x 10-5 (D) 1.8 x 10-5 (Ε) 1.8 x 10-10 53. Which statement below is inconsistent with the concept of isotopes? (Α) Each element is composed of atoms (B) All atoms of an element are identical (C) The atoms of different elements have different chemical and physical properties (D) The combining of elements leads to the formation of compounds (Ε) In a compound, the kinds and numbers of atoms are constant 54. Which sample below has its atoms arranged in a regular, geometric pattern? (Α) NaC2H3O2(s) (B) H2O(l) (C) Ar(g) (D) NaCl(aq) (Ε) CH4(g) 55. Of the statements below, which holds true for the elements found in Na2HPO4? (Α) The total molar mass of 71 grams/mole (B) The percent by mass of oxygen is 45% (C) The percent by mass of sodium is 16% (D) The percent by mass of phosphorus is 44% (Ε) The percent by mass of hydrogen is 13% 56. Carbon and oxygen react to form carbon dioxide according to the reaction: C(s) + O 2(g) CO2(g). How much carbon dioxide can be formed from the reaction of 36 grams of carbon with 64 grams of oxygen gas? SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 128

(Α) 36 grams

(B) 64 grams

(D) 132 grams (Ε) 88 grams

57. What is the correct mass-action expression for the reaction: 2NO(g) + Cl2(g) ⇌ 2NOCl(g)? (Α)

(B)

(C)

(Ε)

58. Which of the following processes will decrease the rate of a chemical reaction? i. Using highly concentrated reactants ii. Decreasing the temperature by 25 K iii. Stirring the reactants (Α) i only (B) ii only (C) i and iii only (D) ii and iii only (Ε) i, ii, and iii 59. Of the substances below, which is best able to conduct electricity? (Α) KBr(l) (B) NaC2H3O2(s) (C) C6H12O6(aq) (D) CH3OH(aq) (Ε) NaCl(s) 60. A voltaic cell is set up and a chemical reaction proceeds spontaneously. Which of the following will not occur in this reaction? (Α) The electrons will migrate through the wire (B) The cations in the salt bridge will migrate to the anode half-cell (C) The cathode will gain mass (D) The anode will lose mass (Ε) Reduction will occur at the cathode 61. What is the value for ΔH for the reaction: D + A + B F A + B C ΔH = -390 kJ D + ½B E ΔH = -280 kJ F + ½B C + E ΔH = -275 kJ (Α) -165 kJ (B) +385 kJ (C) -395 kJ (D) -945 kJ

(Ε) +400 kJ

62. The oxidation state of the elements in the choices below will be -1 except for: (Α) F in HF (B) Cl in NaCl (C) O in H2O2 (D) F in NaF (Ε) H in Na2HPO4 63. Which substance is not correctly paired with the type of bonding found between the atoms of that substance? (Α) CH4 - covalent bonds (B) CaO - ionic bonds (C) Fe - metallic bonds (D) H3O+ - coordinate covalent bonds (Ε) Cl2 - polar covalent bonds 64. Which electron configuration shows that of an excited atom? (Α) 1s22s22p63s1 (B) 1s22s22p63s23p63d1 (C) 1s22s22p4 (D) 1s22s22p63s23p64s2 (Ε) 1s22s22p63s23p3 65. Given the chemical reaction: 3H2(g) + N2(g) ⇌ 2NH3(g) + energy, the forward reaction can best be described as a(n): SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 129

i. Synthesis reaction ii. Phase equilibrium (Α) ii only (B) i and ii only (C) i and iii only (Ε) i, ii and iii

iii. Exothermic reaction (D) ii and iii only

66. Which of the following is not true regarding conjugates and conjugate pairs? (Α) HF and F- are conjugate pairs (B) NaC2H3O2 and C2H3O2- are conjugate pairs (C) CO32- is the conjugate base of HCO3- (D) NH4+ is the conjugate acid of NH3 (Ε) A conjugate pair will differ by an H+ ion 67. What is the ratio of the rate of effusion of hydrogen gas to that of helium gas? (Α) 1.41 (B) 2.00 (C) 4.00 (D) 0.50 (Ε) 1.00 68. Which substance below will exhibit hydrogen bonding between the molecules of the substance? (Α) CH4 (B) HBr (C) HCl (D) H2O (Ε) H2 69. Given the reaction: N2(g) + 3H2(g) ⇌ 2NH3(g) + 22kcal, what is the value of ΔH for the reverse reaction when 6 moles of NH3 are consumed to produce nitrogen gas and hydrogen gas? (Α) +22 kcal (B) +66 kcal (C) -22 kcal (D) -66 kcal (Ε) +33 kcal 70. A titration is set up so that 40.0 mL of 1.0 M NaOH are titrated with 2.0 M HCl. If the initial reading of the meniscus of the acid’s buret is 3.15 mL, what will the final buret reading be? (Α) 20.00 mL (B) 40.00 mL (C) 43.15 mL (D) 23.15 mL (Ε) 13.15 mL 71. Which of the following best describes the orbital overlap in a molecule of:

i. s to s (Α) i only (Ε) i, ii and iii

ii. s to p (B) ii only

iii. sp2 to sp2 (C) i and iii only

(D) ii and iii only

72. What will be the change in the freezing point of the water in a solution of 1 m NaCl(aq)? (Α) -1.86oC (B) -0.52oC (C) -3.72oC (D) 1.86oC (Ε) 3.72oC 73. Which metal will not generate hydrogen gas when placed in HCl(aq)? (Α) Au (B) Mg (C) Ca (D) Sr (Ε) Zn 74. Which substance is the best oxidizing agent? (Α) Fe (B) O2 (C) Na (D) Li

(Ε) F2

75. Which substance is not correctly paired with the bonding found between the molecules of that substance? SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 130

(Α) NH3—hydrogen bonding (B) F2—Van der Waals (dispersion) forces (C) HCl—dipoles (D) CH4—dipoles (Ε) NaCl(aq)—molecule-ion attraction 76. Which solution is not expected to conduct electricity? (Α) NaCl(aq) (B) C6H12O6(aq) (C) KBr(aq) (D) HC2H3O2(aq) (Ε) NaOH(aq) 77. Which of the following statements about solubility is correct? (Α) Gases decrease in solubility with an increase in temperature (B) NaCl is insoluble in water (C) PbI2 is soluble in water (D) All nitrates are insoluble in water (Ε) Solubility depends solely upon the amount of solvent used 78. In 6.20 hours, a 50.0-gram sample of 112Ag decays to 12.5 grams. What is the half-life of 112Ag? (Α) 1.60 hours (B) 3.10 hours (C) 6.20 hours (D) 12.4 hours (Ε) 18.6 hours 79. The modern periodic table is based upon: (Α) Atomic mass of the elements (B) Number of neutrons in the nucleus (C) Number of isotopes of an element (D) Oxidation states (Ε) Number of protons in the nucleus 80. The prefix centi- means: (Α) One thousand (B) One thousandth (D) One hundredth (Ε) One millionth

c. One hundred

81. What is the pH of a 0.1 M acid solution where the acid has a Ka of 1 x 10-5? (Α) 3 (B) 5 (C) 6 (D) 4 (Ε) 1 82. Which of the following would you not do in a laboratory setting? i. Pour acids and bases over a sink ii. Wear goggles iii. Heat a stoppered test tube (Α) i only (B) ii only (C) iii only (D) i and iii only

(Ε) i, ii, and iii

83. Which of the following statements is not true regarding the kinetic molecular theory? (Α) The volume that gas molecules occupy is negligible compared to the volume within which the gas is contained: (B) There are no forces present between gas molecules (C) Collisions between gas molecules are perfectly elastic d. Gas molecules travel in a continuous, random motion (Ε) The average kinetic energy of gas molecules is inversely proportional to temperature 84. How many times more basic is a solution with a pH of 10 than a solution with a pH of 8: SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 131

(Α) A pH of 10 is two times as basic (C) A pH of 10 is 2,000 times as basic (Ε) A pH of 10 is 100 times as basic

(B) A pH of 8 is two times as basic (D) A pH of 8 is 20 times as basic

85. Which of the following reactions is not labeled correctly? (Α) Fe + Cr3+ Fe3+ + Cr (redox) (B) KBr + H2O (C) CH4 + 2O2 CO2 + 2H2O (combustion) (D) CH4 + Cl2 (Ε) CO2 + H2O H2CO3 (synthesis)

SAT CHEMISTRY Dr. D. Bampilis

HBr + KOH (hydrolysis) CH3Cl + HCl (addition)

Σελίδα 132

Practice test 4

For 1 - 4: (Α) The point of equilibrium (B) The triple point (D) The point where reactants first form products

1. a specific temperature and pressure where solid, liquid, and gas phases exist simultaneously 2. can be shifted by adding more reactants 3. vapor pressure of a liquid is equal to the pressure of the surroundings 4. the activated complex For 5 - 8: (Α) red

(B) purple

(D) green

(Ε) blue

5. copper(II) sulfate solution 6. chlorine gas 7. KMnO4 solution 8. bromine solution For 9 - 11: (Α) voltaic cell (D) pH meter

(B) electrolytic cell (Ε) calorimeter

9. requires an external current to make a redox reaction spontaneous 10. requires a salt bridge 11. detects radioactive particles For 12 - 15: (Α) halogens (D) noble gases

(B) alkali metals (Ε) lanthanides

12. valence electrons are located in the f orbitals 13. need to lose one electron to form a stable octet SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 133

14. will have the highest first ionization energies 15. contain elements in the solid, liquid and gas phases at STP For 16 - 19: (Α) 9.03 x 1023 molecules (D) 6.0 grams

(B) 44.8 liters (Ε) 3.01 x 1023 atoms

16. 0.25 moles of O2 at STP 17. 3.0 moles of H2 at STP 18. 56 grams of N2 at STP 19. 96.0 grams of SO2 at STP For 20 - 22: (Α) Water (Ε) Sodium chloride

(B) Hydrogen bromide

(D) Argon

(B) beta particle (Ε) deuteron

20. hydrogen bonding 21. dipoles 22. dispersion forces For 23 - 25: (Α) alpha particle (D) positron 23. Po-218

At-218 + X

24. Tc-99

Tc-99 + X

25. Ne-19

F-19 + X

Q Statement I BECAUSE Statement II 26. An element’s nuclear charge is equal to the number of protons in the nucleus BECAUSE The only charged particles in the nucleus are neutrons 27. A reaction will be spontaneous if ΔH is negative and ΔS is positive BECAUSE ΔG will be negative when there is a decrease in enthalpy and an increase in entropy

SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 134

28. Cl- is the conjugate base of HCl BECAUSE A conjugate base is formed when an acid gains a proton 29. An electrolytic cell makes a non spontaneous redox reaction occur BECAUSE An electrolytic cell uses an external current to drive a redox reaction 30. The maximum number of electrons allowed in the third principal energy level is 18 BECAUSE The maximum number of electrons allowed in a principal energy level is dictated by the equation 2n2 31. 3000 kilograms is equal to 3 grams BECAUSE The prefix kilo- means “one thousandth” 32. An increase in temperature will cause a gas to expand BECAUSE Temperature and volume have a direct relationship 33. A catalyst will change the heat of reaction BECAUSE A catalyst will lower the potential energy of the activated complex in a reaction 34. Helium will have fewer dispersion forces between its atoms than the other noble gases BECAUSE As the mass of non polar atoms and molecules increases, dispersion forces increase 35. Nitrogen gas will have a greater rate of effusion than oxygen gas BECAUSE Lighter, less dense gases travel faster than heavier, more dense gases 36. Propane can be decomposed chemically BECAUSE Propane is a compound that is made up of simpler elements 37. A mixture of two different liquids can be separated via distillation BECAUSE Different liquids have different boiling points 38. Isotopes have different atomic numbers BECAUSE Isotopes must have different numbers of electrons

SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 135

39. Butene can be converted into butane BECAUSE The addition reaction of hydrogen gas to an alkene will form and alkane 40. NaCl is a basic salt BECAUSE Hydrolysis of NaCl reveals the formation of NaOH and HCl 41. When 58 grams of water is heated from 275 K to 365 K, the water: (Α) Absorbs 21,820 J (B) Absorbs 377 J (C) Releases 5,220 J (D) Absorbs 242 J (Ε) Releases 90 J 42. Which of the following are uses for radiation and radioactivity that are of benefit to us? i. Nuclear waste ii. Radioisotopes iii. Excess exposure (Α) i only (B) ii only (C) iii only (D) i and ii only (Ε) i and iii only 43. Which of the following statements is not part of the kinetic molecular theory? (Α) The average kinetic energy of gas molecules is directly proportional to temperature. (B) Attractive and repulsive forces are present between gas molecules. (C) Collisions between gas molecules are perfectly elastic. (D) Gas molecules travel in a continuous, random motion. (Ε) The volume that gas molecules occupy is minimal compared to the volume within which the gas is contained. 44. The following redox reaction occurs in an acidic solution: Ce4+ + Bi the coefficient before the Ce4+ when the equation is fully balanced? (Α) 1 (B) 2 (C) 3 (D) 6 (Ε) 9

Ce3+ + BiO+. What is

45. Which statement regarding significant figures is false? (Α) Zeros can be significant (B) When multiplying, the answer is determined by the number of significant figures (C) When adding, the answer is determined by the number of decimal places (D) When dividing, the answer is determined by the number of decimal places (Ε) The number 50,004 has five significant figures 46. Which statement below best describes the molecule in question? (Α) Water has a bent molecular geometry and one lone pair of electrons (B) Ammonia has a trigonal pyramidal molecular geometry and two long pairs of electrons (C) Methane has a trigonal planar molecular geometry (D) Carbon dioxide is linear because it has one single bond and one triple bond (Ε) The carbon atoms in ethane are sp3 hybridized 47. A compound was analyzed and found to be 12.1% C, 71.7% Cl, and 16.2% O. What is the empirical formula for this compound? (Α) C2OCl (B) COCl (C) CO2Cl2 (D) C2O2Cl (Ε) CCl2O SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 136

48. Which statement is true about the percent composition by mass of C6H12O6? (Α) Carbon is 6.7% by mass (B) Oxygen is 53.3% by mass (C)Hydrogen is 12% by mass (D) Carbon is 72% by mass (Ε) Carbon is 20% by mass 49. Which process would have a positive value for the change in entropy? i. The expansion of the universe ii. The condensation of a liquid iii. A food fight in a school cafeteria (Α) i only (B) ii only (C) iii only (D) ii and iii only (Ε) i and iii only 50. Of the gases below, which would react with rain water to produce acid rain? i. CFCs ii. Methane iii. Carbon dioxide (Α) i only (B) ii only (C) iii only (D) i and iii only (Ε) i, ii and iii 51. A sample of gas is trapped in a manometer and the stopcock is opened. The level of mercury moves to a new height as can be seen in the diagram. If the pressure of the gas inside the manometer is 815 torr, what is the atmospheric pressure in this case? (Α) 760 torr (B) 740 torr (C) 750 torr (D) 815 torr (Ε) 880 torr 52. Which aqueous solution is expected to have the highest boiling point? (Α) 1.5 m FeCl2 (B) 3.0 m CH3OH (C) 2.5 m C6H12O6 (D) 2.5 m NaCl (Ε) 1.0 m CaCl2 53. Which Ka value is that of a better electrolyte? (Α) 1.0 x 10-2 (B) 2.0 x 10-12 (C) 5.0 x 10-7 (Ε) 1.0 x 10-6

(D) 3.0 x 10-4

54. The following substances were all dissolved in 100 grams of water at 290 K to produce saturated solutions. If the solution is heated to 310 K, which substance will have a decrease in its solubility? (Α) NaCl (B) KI (C) CaCl2 (D) HCl (Ε) KNO3 55. Methane undergoes a combustion reaction according to the reaction CH4(g) + 2O2(g) CO2(g) + 2H2O(l). How many grams of methane gas were burned if 67.2 liters of carbon dioxide gas are produced in the reaction? (assume STP) (Α) 16 grams (B) 48 grams (C) 3 grams (D) 132 grams (Ε) 22.4 grams 56. A closed system contains the following reaction at STP: Cl2(g) + 2NO2(g) ⇌ 2NO2Cl(g). What is the equilibrium constant expression for this reaction? (Α)

(B)

(C)

(D)

(Ε)

57. At a particular temperature, the equilibrium concentrations of the substances in the previous question are as follows: [NO2Cl] = 0.5 M, [Cl2] = 0.3 M, [NO2] = 0.2 M. What is the value of the equilibrium constant for this reaction? SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 137

(Α) 2.1

(B) 0.48

(D) 20.83

(Ε) 208.83

58. Which Lewis structure below has been drawn incorrectly? (Α) H:H

(B) H:C:::N:

(C)

(D):N:::N:

(Ε) F: B : F F 59. Which reaction below demonstrates the Lewis definition of acids and bases? (Α) HCl + NaOH HOH + NaCl (B) H2O + NH3 OH- + NH4+ (C) NH3 + BF3 NH3BF3 (D) HI + KOH H2O + KI (Ε) H+ + OH H2O 60. Which sample is a homogeneous mixture? (Α) KI(aq) (B) Fe(s) (C) CO2(g) (D) NH3(l)

(Ε) NaCl(s)

61. Which pair below represents isomers of the same compound? (Α) CH3CH2CH2OH and HOCH2CH2CH3 (B) CH3CH2CH3 and CH3CH2CH2CH3 (C) CH3CH(Cl)CH3 and CH3CH2CH2Cl (D) CH3COCH3 and CH3CH2CH2CHO (Ε) ClCH2CH2Br and BrCH2CH2Cl 62. Which would you never to in a laboratory setting? i. Eat and drink in the laboratory ii. Push a thermometer through a rubber stopper iii. Remove your goggles to take a better look at a reaction (Α) i only (B) ii only (C) iii only (D) i and iii only

(Ε) i, ii and iii

63. How many pi bonds are there in a molecule of N C-CH2-CH2-CO-NH-CH=CH2 (Α) 7 (B) 4 (C) 12 d. 10 (Ε) 5 64. When the equation: C2H6 + O2 CO2 + H2O is completely balanced using the lowest whole number coefficients, the sum of the coefficients will be: (Α) 4 (B) 9.5 (C) 19 (D) 15.5 (Ε) 11 65. From the heats of reaction of these individual reactions: A + B 2C ΔH = -500 kJ D + 2B E ΔH = -700 kJ 2D + 2A F ΔH = +50 kJ Find the heat of reaction for F + 6B 2E + 4C (Α) +450 kJ (B) -1100 kJ (C) +2350 kJ (D) -350 kJ

(Ε) -2450 kJ

66. Which solutions have a concentration of 1.0 M? i. 74 grams of calcium hydroxide dissolved to make 1 liter of solution ii. 74.5 grams of potassium chloride dissolved to make 1 liter of solution SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 138

iii. 87 grams of lithium bromide dissolved to make 1 liter of solution (Α) i only (B) iii only (C) i and iii only (D) ii and iii only (Ε) i, ii and iii 67. According to the reaction 3H2 + N2 2NH3, how many grams of hydrogen gas and nitrogen gas are needed to make exactly 68 grams of ammonia? (Α) 2 grams of hydrogen gas and 28 grams of nitrogen gas (B) 3 grams of hydrogen gas and 1 gram of nitrogen gas (C) 12 grams of hydrogen gas and 56 grams of nitrogen gas (D) 102 grams of hydrogen gas and 34 grams of nitrogen gas (Ε) 6 grams of hydrogen gas and 2 grams of nitrogen gas 68. Which compound is not paired with its correct name? (Α) FeCl2 / iron(II) chloride (B) K2O / potassium oxide c. NO2 / nitrogen dioxide (D) PCl3 / potassium trichloride (Ε) NH4Cl / ammonium chloride 69. How many grams of HI can be made from 6 grams of H2 and 800 grams of I2 in the following reaction: H2 + I2 2HI? (Α) 800 grams of HI can be made with 38 grams of excess iodine (B) 768 grams of HI can be made with 6 grams of excess hydrogen (C) 768 grams of HI can be made with 38 grams of excess iodine (D) 2286 grams of HI can be made with no excess reactants (Ε) 806 grams of HI can be made with no excess reactants 70. 500 mL of a 0.2 M solution has 200 mL of water added to it. What is the new Molarity of this solution? (Α) 0.50 M (B) 0.28 M (C) 0.70 M (D) 0.14 M (Ε) 0.40 M 71. Which mixture is correctly paired with a method for separation of the mixture? (Α) Oil and water - filter paper (B) Salt water - distillation (C) Sand and water - separatory funnel (D) Sand and sugar - tweezers (Ε) Sugar water - filter paper 72. Which reaction between ions does not form a precipitate? (Α) Ag+ and Cl(B) Pb2+ and 2I(C) Ca2+ and CO32(D) Hg2+ and 2Br(Ε) Na+ and OH73. Which will happen when sodium sulfate is added to a saturated solution of CaSO4 that is at equilibrium? [CaSO4(s) ⇌ Ca2+(aq) + SO42-(aq)] (Α) The solubility of the calcium sulfate will decrease (B) The concentration of calcium ions will increase (C) The reaction will shift to the right (D) The Ksp value will change (Ε) The equilibrium will shift to consume the decrease in sulfate ions.

SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 139

74. Given the reaction: 2A(g) + B(g) + heat ⇌ 3C(g) + D(g), what could be done to the reaction to shift the equilibrium so that more D is made? (Α) Increase the concentration of D (B) Increase the concentration of C (C) Increase the temperature (D) Increase the pressure (Ε) Remove B from the reaction 75. A 16-gram sample of water at 273 K is cooled so that it becomes a completely solid ice cube at 273 K. How much heat was released by the sample of water to form this ice cube? (Α) 16 J (B) 4368 J (C) 18258 J (D) 350 J (Ε) 5334 J 76. Sublimation is the process by which a solid becomes a gas without having a liquid phase. Which of these substances can sublime? i. Iodine ii. Naphthalene iii. Carbon dioxide (Α) i only (B) ii only (C) iii only (D) i and iii only (Ε) i, ii and iii 77. Which of the following will decrease the rate of a reaction? (Α) Using powdered solids instead of whole pieces (B) Selecting ionic reactants that have been dissolved in water (C) Decreasing the temperature (D) Increasing the pressure (Ε) Adding a catalyst 78. Three gases are mixed in a sealed container. The container has 0.3 moles of gas A, 0.4 moles of gas B, and 0.3 moles of gas C. The total pressure of the gases is 660 torr. What is true about the partial pressures of the gases? (Α) The partial pressure of gas A is 264 torr (B) The partial pressure of gas B is 396 torr (C) The partial pressure of gas C is 220 torr (D) The partial pressures of gases A and C are each 198 torr (Ε) The partial pressure of gas B is 660 torr Nickel and aluminum electrodes are used to build a galvanic cell. The standard reduction potential for the nickel(II) ion is -0.26 V and that of the aluminum(III) ion is -1.66 V 79. What is the half reaction that occurs at the cathode? (Α) Al Al3+ + 3e(B) Ni2+ + 2e- Ni (C) Ni Ni2+ + 2e(D) 2Al3+ + 6e- 2Al (Ε) Al3+ + 3e- Al 80. Which statement is true about the setup shown? (Α) The electrode potential for this cell is 1.40 V (B) The electrode potential for this cell is 2.54 V (C) Electrons will be carried by the salt bridge (D) Ions will be carried through the wire (Ε) The reaction is nonspontaneous

SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 140

81. Over a number of years the average pH of a stream changes from a pH of 6.9 to a pH of 9 due to acid rain. Which statement is true about the pH of the stream? (Α) The pH of the stream now is one times more acidic than it was years ago (B) The stream now has 10 times more hydroxide ions than it did years ago (C) The pH of the stream is now 10 times more acidic than it was years ago (D) The stream is more basic now than it was years ago (Ε) The concentration of hydronium ion in the stream has decreased over the years 82. An alkaline earth metal, element M, reacts with oxygen. What is going to be the general formula for the compound formed? (Α) M2O (B) MO (C) MO2 (D) M2O3 (Ε) M3O2 83. Which functional group below does not contain a carbonyl group? (Α) Aldehydes (B) Ketones (C) Esters (D) Ethers (Ε) Carboxylic acids 84. Using the bond dissociation energies found at the end, calculate the change in the heat of reaction for 2H2 + O2 2H2O. (Α) -118 kJ (B) +118 kJ (C) -91 kJ (D) -1042 kJ (Ε) -833 kJ 85. Equilibrium: (Α) Is defined as equal concentrations of reactants and products (B) Is defined as equal rates for forward and reverse reactions (C) Can be shifted by adding a catalyst (D) Can exist for chemical changes but not for physical changes (Ε) Must always favor the formation of products Bond energies: Bond kJ/mol C - C: 349 C - Cl: 329 C - H: 412 C = O: 798 Cl - Cl: 240 H - Cl: 430 H - H: 435 N - H: 390 N - N: 163 N N: 941 O - H: 462 O - O: 145

SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 141

SAT TIPS 1. TEAR OUT THE PERIODIC TABLE. Yes, you read correctly, tear that thing out straight away, it'll save you up to 5 precious minutes of time and reduce the possibility of confusion. There is always a 1-2 minute delay between when the supervisor tells you to turn the page to the beginning of the exam and when you actually start the test (this is so that everyone has gotten to the right page; if you've taken collegeboard tests before you should be aware of this). During that time period *BEGIN* (don't do it completely or else you'll see the next page==if you're caught then you might be wrongly accused of cheating) tearing out the table. Note:The periodic table is on the back of the instruction page, in case you didn't know. 2. Before beginning the test, draw on your table those "arrows" that you learned in chemistry. Draw the arrows showing trends such as electronegativity, atomic radius, ionization energy etc. and remember the exceptions to these trends noble gases - electronegativity 2-13 group 15-16 group –ionization energy Cl , F /S,O – electron affinity. 3. In the exam you always have like 3-5 questions and an answer bank of 5-7 (around that number) of solutions. USUALLY you will not use an answer choice more than once, so if you find yourself choosing the same answer for two questions double check that you're 100% certain. If you're not so sure about one of the two questions then there is a good chance that it is incorrect. Note: This is a theory of mine an not necessarily true in all cases, but has stopped me from making mistakes many times. Also, this is true most often when you have 3 questions and works less often when there are 5 questions.

SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 142

4. If you get to a question where there are two possible answers, choose the less controversial answer (the 'best answer'). For example on the last test I was asked: Which of the following is found in nature as an element? Answer choices: Au, Fe...etc.) I quickly realized that the only two possibilities were iron and gold. But, which one is the better answer? Fe is more reactive and is usually found in compounds so Au was the right answer. On such questions go with your gut feeling and don't double question yourself. Your instinct is usually right. 5. Aim to get all the questions right. Chem's curve is not really generous and if you start telling yourself subconsciously that you can get a couple wrong and still get 800 then you will probably make more mistakes. Remember, only those that miss between 0-4 (sometimes 3) RAW marks or 0-3 questions (sometimes 2) get 800s. Thats not a lot of leeway if you ask me, especially if you compare it to physics (-11 questions=800). True/False section If you know one of the statements is definitely true but you have no clue about the other statement there are a few ways to increase your chances of getting it right. 1. In the T/F section, only worry about the “because” part of the question if both answers are true. In other words, do not worry about whether the second part is the correct explanation (T T CE) if the second part were not true. Just bubble in T F. Otherwise you are wasting time. 2.If you think KNOW that if the second statement, if true, will justify the first statement (T,T CE), then it is most likely that the answer is T T CE. However, if you know the second statement if true, will NOT justify the first statement (T T) then it is more likely that the second statement is false (T F). This works really well on the hard questions which you have no clue about. SAT CHEMISTRY Dr. D. Bampilis

Σελίδα 143

3. If you cannot apply tip 1 to the question, there is still some hope. Statistically, if you only know the first statement is definitely true but you have absolutely no clue about the second statement validity and/or you don't even understand what its stating (it happens sometimes ) then you'll have a greater chance of getting the answer correct if you choose (T,F). There is a simple explanation: So if first one is definitely true then probability of (T,F) is 50%, (T,T) is 25% and (T,T CE) is 25%. 4. You'll usually have at least 2 (T,T CE) answers in any given test. If you have 0 or 5 then something is most likely wrong.

Somethings you should learn: 1. Vapor pressure of water must be subtracted when measuring pressure of a gas collected over water. This is due to Dalton’s Law of Partial Pressures. 2. Learn how to use millimeters of mercury and torr to measure pressure. Just knowing "atmospheres" (e.g. 1 atm) is not enough. 3. Memorize absolute zero (-273 K) 4. Solubility rules: Most alkali metal compounds and NH4+ compounds are soluble. Cl-, Br-, I- compounds are soluble, except when they contain Ag+, Hg22+, or Pb2+. F- compounds are soluble, except when they contain group 2A metals. NO3-, ClO3-, ClO4- and CH3COO- compounds are soluble. SO42- compounds are soluble, except when they include Ca2+, Sr2+, Ba2+, Ag+, Pb2+, or Hg22+. CO32-, PO43-, C2O42-, CrO42-, S2-, OH-, and O2- compounds are insoluble. Group 2A metal oxides are classified as strong bases even though they are not very soluble. Oxides (except group 1A) are insoluble, and when reacted with water, they form either acids (nonmetal oxides) or bases (metal oxides). 5. Learn solubility product ksp Solubility products Ksp - Qsp Ksp - S AgCl(s) AgSO4(s)

Ag+(aq) + Cl-(aq) Ksp = s2

s = K sp

2Ag+(aq) + SO42-(aq) Ksp = 4s3

SAT CHEMISTRY Dr. D. Bampilis

Ksp 4

Σελίδα 144

V k T Boyle’s law P · V = k P k Gay - Lussac’s law T Dalton’s law of partial pressures : Ptotal = P1 + P2 + P3 + … 7. Mendeleev-periodic table Faraday-electrolytic cells Millikan-charge of electrons Rutherford- nucleus Einstein-photoelectric effect Curie- radioactivity 6. Charles’s law

8. Memorize the nomenclature of the relevant organic chemistry 9. Always skip stoichiometry questions (until later). It is the easiest content wise but balancing equations take the longest time. 10. Be careful of ions. Remember that a Calcium 2+ ion has the same number of electrons as Argon. 11. Quantum numbers Principal quantum number (n) Second quantum number or azimuthal quantum number (l ) Magnetic quantum number (ml) Spin quantum number (ms)

SAT CHEMISTRY Dr. D. Bampilis

Has positive values of 1, 2, 3, etc. As n increases, the orbital becomes larger-this means that the electron has a higher energy level and is less tightly bound to the nucleus. Has values from 0 to n - 1. This defines the shape of the orbital, and the value of l is designated by the letters s, p, d, and f, which correspond to values for l of 0, 1, 2, and 3. In other words, if the value of l is 0, it is expressed as s; if l = 1 = p, l = 2 = d, and l = 3 = f. Determines the orientation of the orbital in space relative to the other orbitals in the atom. This quantum number has values from l through 0 to +l. Specifies the value for the spin and is either +1/2 or -1/2. No more than two electrons can occupy any one orbital. In order for two electrons to occupy the same orbital, they must have opposite spins.

Σελίδα 145

12. MEMORIZE tests for different chemicals. Very important to know these Colors of Salts: Salt

Formula

Solid

Hydrated copper sulfate Anhydrous copper sulfate Copper nitrate Copper chloride Copper carbonate Copper oxide Iron(II) salts

CuSO4.5H2O

Blue crystals

CuSO4

White powder

Blue

Cu(NO3)2

Blue

CuCl2

Blue crystals Green

CuCO3

Green

Insoluble

CuO E.g.: FeSO4, Fe(NO3)2 E.g.: Fe(NO3)3

Black Pale green crystals Reddish brown

Insoluble Pale green Reddish brown

Iron(III) salts

In Solution Blue

Green

Tests for Gases: Gas Ammonia Carbon dioxide Oxygen Hydrogen Chlorine Nitrogen dioxide Sulfur dioxide

SAT CHEMISTRY Dr. D. Bampilis

Formula NH3 CO2

Tests Turns damp red litmus paper blue Turns limewater milky

O2 H2 Cl2 NO2

Relights a glowing splint ‘Pops’ with a lighted splint Bleaches damp litmus paper Turns damp blue litmus paper red

SO2

Turns acidified aqueous potassium dichromate(VI) from orange to green

Σελίδα 146

Tests for Anions: Anion Carbonate (CO32-)

Test Add dilute acid

Chloride (Cl-)(in solution)

Acidify with dilute nitric acid, then add aqueous silver nitrate Acidify with dilute nitric acid, then add aqueous silver nitrate Add aqueous sodium hydroxide, then aluminium foil; warm carefully Acidify, then add aqueous barium nitrate

Iodide (I-)(in solution) Nitrate (NO3-)(in solution) Sulfate (SO42-)

Result Effervescence, carbon dioxide produced White ppt.

Yellow ppt.

Ammonia produced

White ppt.

Tests for aqueous cations: Cation Aluminium (Al3+) Ammonium (NH4+) Calcium (Ca2+) Copper (Cu2+)

Iron(II) (Fe2+) Iron(III) (Fe3+) Zinc (Zn2+)

Effect of aqueous sodium hydroxide White ppt., soluble in excess giving a colourless solution Ammonia produced on warming White ppt., insoluble in excess Light blue ppt., insoluble in excess

Green ppt., insoluble in excess Red-brown ppt., insoluble in excess White ppt., soluble in excess, giving a colourless solution

SAT CHEMISTRY Dr. D. Bampilis

Effect of aqueous ammonia White ppt., insoluble in excess – No ppt. or very slight white ppt. Light blue ppt., soluble in excess, giving a dark blue solution Green ppt., insoluble in excess Red-brown ppt., insoluble in excess White ppt., soluble in excess, giving a colourless solution

Σελίδα 147

Flame colors and the alkali metal ion they represent Flame colour

Ion present

Red

Lithium, Li+

Orange

Sodium, Na+

Lilac

Potassium , K+

Brick red

Calcium, Ca2+

13. Graham’s law of diffusion (effusion)

Rate A Rate B

MB MA

14. Mercury Barometer

15. Manometer

SAT CHEMISTRY Dr. D. Bampilis

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16. Phase diagram

17. Formation of gases Gases: HX, H2S, HCN, SO2(H2SO3), CO2(H2CO3), NH3(NH4OH) 18. Learn how to balance redox reactions. This is much more difficult than simple stoichiometry. 19. Strong acids bases There are six strong acids that completely ionize: HCl, HBr, HI, HNO3, H2SO4, HClO4. All other acids are weak and are written together, as molecules. The strong bases that ionize are oxides and hydroxides of group 1A and 2A metals. All other oxides and hydroxides are considered weak and written together, as molecules. 20. Thermochemistry Calorimetry: q = mcΔΤ q = mL (phase change) Hess’s law o o rxn f (products) f (reac tan ts) rxn

o c

(reac tan ts)

o c

(products)

Bond dissociation energy (bond length) SAT CHEMISTRY Dr. D. Bampilis

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Bond broken: Energy absorbed Bond formed: Energy evolved o D broken Dformed rxn Entropy – ΔS > 0 • solid → liquid gas •T • mixing different particles •

n (gas) Rxn ΔG = -

ΔΗ -

ΤΔS

+ +

+ -

Spontaneous spontaneous at low T Nonspontaneous spontaneous at high T

21. Freezing Point Depression DT f = K f x m solute x i where: DTf = the change in freezing point Kf = molal freezing point depression constant for the substance (for water = 1.86oC/m) m = molality of the solution i = number of ions in solution (this is equal to 1 for covalent compounds and is equal to the number of ions in solution for ionic compounds) 22. Boiling Point Elevation DT b = K b x m solute x i where: Kb = molal boiling point elevation constant (for water = 0.51oC/m) 23. Vapour pressure of solution P N Po N n 24. Osmotic Pressure Π V = i n RT

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26. Solution Colloid Suspension

Solution

Suspension Colloid Cloudy, heterogeneous, at Cloudy but uniform and least two substances homogeneous visible

Appearance

Clear, transparent and homogeneous

Particle Size

molecule in size

Effect of Light Tyndall Effect

none -- light passes through, particles do variable not reflect light

Effect of Sedimentation

none

larger than 10,000 Angstroms

particles will eventually settle out

10-1000 Angstroms

light is dispersed by colloidal particles

none

27. Calculating an Equilibrium Constant from the Free Energy Change If we know the standard state free energy change, Go, for a chemical process at some temperature T, we can calculate the equilibrium constant for the process at that temperature using the relationship between Go and K.

Rearrangement gives

At standard conditions, DG° = - n F DE° 28. Normality Equivalence defined through the amount of substance which will either: react with or supply one mole of hydrogen ions (H+) in an acid–base reaction; or react with or supply one mole of electrons in a redox reaction 29. Nuclear chemistry Alpha particles a or 42 He 2

( 42 He)

Beta particles β- or o1 e or Gamma Rays γ High energy radiation SAT CHEMISTRY Dr. D. Bampilis

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Control Z – A • a Zdaughter Zparent

A daughter • β Zdaughter • γ

A parent

Zparent 1

A daughter

A parent

Zdaughter

Zparent

A daughter

A parent

1 Fraction of original nuclei remaining after n half – lives: ( ) n 2

1 Fraction of nuclei that has decayed away after n half – lives: 1 ( ) n 2 30. Common alloys Alloy Bronze Brass

Composition 90% Cu 10% Sn 67% Cu 33% Zn

Properties Strong, corrosion resistant, easily cast

Uses Sculptures, church bells

bright gold appearance, higher malleability than bronze or zinc

musical instruments

Sterling 93% Ag silver 7% Cu Duralumin 95% Al 4% Cu 1% Mn and Mg

Harder than pure silver, unreactive, lustrous Low density, very strong

Jewellery

Stainless steel

Resistant to corrosion,

Cutlery and surgical instruments

Hard, tough, and corrosionresistant

Bridges , buildings

Steel Pewter

Solder

Amalgams

74% Fe 18% Cr, 8% Ni 99% Fe 0.3–0.6% C 85–99% Sn 1–2% Cu 1.5-6% Sb 70% Pb 30%Sn 67% - 74% Ag 25-28% Sn, up to 6% Cu, 2% Zn 3% Hg

SAT CHEMISTRY Dr. D. Bampilis

Low melting point and stronger than lead or tin.

Aircraft bodies, racing bikes

Dishes, church vessels, and decorative items Connections in electr ical wiring and in plumbing Teeth fillings

Σελίδα 152

ANSWERS 1. Atomic structure 1. A 2. D 3. B 4. A 5. E 6. C 7. A 8. D 9. D 10. A 11. B 12. D 13. D

2. Periodicity 1. E 2. D 3. A 4. D 5. C 6. T - T 7. T - F 8. T - F 9. T – T - CE 10. T-T

Atomic structure & periodicity 1. C 2. D 3. A 4. D 5. A 6. C 7. B 8. A 9. D 10. C 11. D 12. E 13. D 14. A 15. B 16. A 17. C 18. TF 19. T T CE 20. T F 21. F T 22. F T 23. T F 24. F F 25. T T CE 26. T F 27. T T CE 28. T T 29. T T CE 30. T T CE 31. F F 32. T F 33. C 34. B 35. C 36. E 37. E 38. B 39. C 40. A 41. A 42. E 43. C 44. A 45. D 46. C 47. E 48. A 49. C 50. B 51. D 52. C 53. A 54. A 55. D 56. C 57. C 58. D 59.Β 60.Β

3. Bonding 1. A 2. B 3. C 4. E 5. B 6. D 7. C 8. E 9. C 10. D 11. B 12. D 13. B 14. A 15. C 16. D 17. A 18. D 19. B 20. C 21. D 22. A 23. T, T, CE 24. T, T, CE 25. T, T, CE 26. F, T 27. T, T, CE 28. F, T 29. T, F 30. T, T 31. T, T, CE 32. T, T, CE 33. D 34. B 35. D 36. C 37. B 38. E 39. B 40. D 41. B 42. A 43. D 44. C 45. D 46. A 47.C 48.B 49.C 50.T – T - CE 51.T - T 52.F - T 53.T - F 54.T – T - CE

4. Gases and Gas Laws 1. D 2. D 3. C 4. C 3. B 6. E 7. C 8. F - T

5. Stoichiometry 1. E 2. B 3. A 4. D 5. B 6. E 7. D 8. B 9. A 10. E 11. C 12. D 13. E 14. B 15. C 16. C 17. E 18. C 19. A 20. C 21. E 22. D 23. D 24. C 25. D 26. A 27. C 28. B 29. C 30. E 31. A 32. C 33. E 34. B 35. C 36. E 37. D 38. C 39. D 40. C 41. D 42. C 43. B 44. B 45. E 46. D 47. A 48. C 49. A 50. A 51. A 52. C 53. C 54. D 55. B 56. A 57. E 58. B 59. D 60. C 61.D 62. D 63. B 64. B 65. F - T

6. Solids, Liquids and Phase Changes 1. E 2. A 3. B 4. C 5. D 6. B 7. A 8. B 9. C 10. D 11. T, T, CE 12. T, F 13. F, T 14. T, T, CE 15. T, T, CE 16. T, F 17. F, T 18. F, F 19. T, T, CE 20. B 21. B 22. A 23. B 24. C 25. C 26. C 27. D 28. A 29. B 30. D 31. B 32. E 33. C 34. A 35. A 36. B 37. C 38. C 39. E 40. C 41. B 42. E 43. A 44. B 45. C 46. B 47. C 48. A 49. D 50. B 51. A 52. A 53. B 54. C 55. D 56. D 57. A 58. C 59.A 60. A 61. A 62. A 63. D 64. A 65. A 66. Α 67. B 68. A 69. 70. A 71.D 72. A 73. A 74. E 75. C 76. A 77. C 78.D 79. D 80. C 81. D 82. E 83. C 84. B 85. T - F 86. T – T - CE 87. T – T - CE 88. C 89. A

7. Reaction Types 1. B 2. E 3. D 4. E 5. E 6. B 7. D 8. A 9. B 10. A 11. D 12. B 13. A SAT CHEMISTRY Dr. D. Bampilis

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8. Thermodynamics 1. B 2. C 3. A 4. A 5. D 6. A 7. D 8. C 9. C 10. A 11. C 12. E 13. D 14. F, T 15. T, T, CE 16. T, T 17. F, F 18. F, F 19. F, F 20. F, T 21. D 22. A 23. A 24. E 25. C 26. D 27. C 28. B 29. E 30. D 31. A 32. E 33. D 34. C 35. C 36. C 37. C 38. C 39. B 40. A 41. E 42. T - T - CE

9. Kinetics 1. A 2. D 3. C 4. E 5. A 6. D 7. B 8. E 9. C 10. A 11. A 12. B 13. E 14. F, F 15. T, T, CE 16. F, T 17. F, F 18. F, T 19. F, T 20. T, F 21. T, T 22. T, T 23. F, T 24. F, F 25. T, T 26. F, T 27. C 28. C 29. D 30. A 31. D 32. B 33. C 34. B 35. A 36. A 37. C 38. T - T 39. C 40. C 41. C 42. E 43. E

10. Equilibrium 1. D 2. C 3. A 4. D 5. D 6. A 7. B 8. E 9. D 10.C 11. D 12. A 13. Ε 14.T - F 15. T, T, CE 16. T, T, CE 17. E 18. C 19. C 20. D 21. A 22. E 23. A 24. C 25. D 26. C

11.Solutions 1. D 2. A 3. A 4. D 5. B 6. E 7. D 8. C 9. B 10. A 11. B 12. T F 13. T T CE 14. F T 15. T F 16. T T CE 17. T T CE 18. B 19. B 20. D 21. D 22. C 23. D 24. D 25. B 26. C 27. C 28. C 29. B 30. C 31. E 32. B 33. E 34. B 35. B 36. E 37. C 38. C 39. B 40. A 41. C 42. B 43. C 44. B 45. C 46. D 47. B 48. D

12. Acids, Bases, Salts 1. B 2. B 3. C 4. A 5. E 6. C 7. C 8. B 9. D 10. D 11. E 12. B 13. E 14. B 15. C 16. C 17. B 18. E 19. T T CE 20. T T CE 21. F F 22. F T 23. T F 24. F F 25. T T 26. F T 27. T T CE 28. T T 29. F T 30. F T 31. T T CE 32. F T 33. C 34. D 35. C 36. A 37. D 38. B 39. C 40. E 41. D (mixture is called “aqua regia”) 42. B 43. D 44. D 45. D 46. A 47. B 48. D 49. B 50. C 51. C 52. D 53. A 54. D 55. B 56. E 57. E 58. E 59.C 60. E 61. C 62. C 63. E 64. A 65. C 66. B 67. C 68. B

13. Redox and Electrochemistry 1. C 2. A 3. D 4. E 5. B 6. A 7. D 8. B 9. B 10. A 11. E 12. A 13. C 14. E 15. B 16. F F 17. F T 18. F T 19. F F 20. T T CE 21. F T 22. F T 23. F T 24. T T CE 25. B 26. D 27. A 28. C 29. C 30. E 31. A 32. D 33. E 34. B 35. E 36. B 37. B 38. B 39. D 40. B 41. C 42. A 43. A 44. D 45. B 46. Ε 47. Α 48.B 49. A 50. A 51. E 52. D 53. C 54. C 55. D 56. C 57.A 58. T – F 59. T - T 60. T – T – CE

14. Organic Chemistry 1. T, T, CE 2. T, T, CE 3. T, F 4. T, T 5. T, T, CE 6. T, T, CE 7. T, T, CE 8. T, T, CE 9. F, F 10. T, T, CE 11. A 12. A 13. C 14. D 15. C 16. B 17. E 18. D 19. E 20. D 21. C 22. C 23. D 24. A 25. D 26. C 27. D 28. E 29. C 30. C 31. D 32. B 33. C 34. B 35. C 36. B 37. E 38. D 39. B 40. A 41. T F

15. Nuclear Chemistry 1. D 2. B 3. C

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16. Descriptive Chemistry 1. B 2. E 3. A 4. B 5. E 6. D 7. T – T – CE

17. Laboratory 1. B 2. C 3. B 4. B 5. A 6. D 7. B 8. A 9. D 10.C 11. E

Practice test 1 1. E 2. C 3. B 4. D 5. B 6. A 7. D 8. E 9. A 10. D 11. B 12. C 13. E 14. D 15. A 16. C 17. A 18. D 19. E 20. D 21. C 22. B 23. A 24. A 25. B 26. T F 27. T T CE 28. F T 29. F T 30. T F 31. T F 32. T T CE 33. T F 34. T T CE 35. T F 36. T T CE 37. F T 38. T T CE 39. F T 40. T T CE 41. B 42. B 43. D 44. E 45. C 46. E 47. B 48. A 49. D 50. A 51. D 52. B 53. C 54. A 55. E 56. D 57. C 58. C 59. E 60. B 61. D 62. A 63. E 64. A 65. E 66. E 67. D 68. B 69. C 70. C 71. D 72. A 73. D 74. C 75. E 76. A 77. B 78. B 79. E 80. A 81. C 82. D 83. C 84. D 85. E

Practice test 2 1. B 2. D 3. C 4. A 5. B 6. C 7. D 8. B 9. C 10. E 11. A 12. E 13. C 14. D 15. E 16. C 17. B 18. A 19. C 20. D 21. B 22. B 23. B 24. C 25. D 26. T T CE 27. F F 28. T F 29. T T CE 30. T T CE 31. T T CE 32. T T CE 33. F F 34. T T CE 35. T T CE 36. T F 37. T F 38. T F 39. F T 40. T F 41. D 42. A 43. B 44. D 45. C 46. B 47. E 48. E 49. A 50. D 51. E 52. C 53. B 54. D 55. A 56. E 57. C 58. C 59. D 60. A 61. B 62. D 63. B 64. E 65. E 66. B 67. B 68. E 69. C 70. A 71. D 72. B 73. C 74. C 75. E 76. D 77. A 78. E 79. D 80. B 81. B 82. D 83. D 84. E 85. C

Practice test 3 1. E 2. B 3. A 4. C 5. A 6. C 7. E 8. B 9. D 10. E 11. B 12. A 13. E 14. E 15. B 16. C 17. B 18. D 19. E 20. A 21. C 22. D 23. A 24. C 25. A 26. T T CE 27. T F 28. F F 29. T F 30. T T CE 31. T T CE 32. F T 33. F T 34. T T CE 35. T F 36. F T 37. T T CE 38. T F 39. T T CE 40. T T CE 41. C 42. D 43. A 44. B 45. D 46. D 47. E 48. A 49. B 50. C 51. B 52. E 53. B 54. A 55. B 56. E 57. D 58. B 59. A 60. B 61. C 62. E 63. E 64. B 65. C 66. B 67. A 68. D 69. B 70. D 71. D 72. C 73. A 74. E 75. D 76. B 77. A 78. B 79. E 80. D 81. A 82. C 83. E 84. E 85. D

Practice test 4 1. B 2. A 3. E 4. D 5. E 6. D 7. B 8. A 9. B 10. A 11. C 12. E 13. B 14. D 15. A 16. E 17. D 18. B 19. A 20. A 21. B 22. D 23. B 24. C 25. D 26. T F 27. T T CE 28. T F 29. T T CE 30. T T CE 31. F F 32. T T CE 33. F T 34. T T CE 35. T T CE 36. T T CE 37. T T CE 38. F F 39. T T CE 40. F T 41. A 42. B 43. B 44. C 45. D 46. E 47. E 48. B 49. E 50. C 51. C 52. D 53. A 54. D 55. B 56. D 57. D 58. E 59. C 60. A 61. C 62. E 63. B 64. C 65. E 66. E 67. C 68. D 69. C 70. D 71. B 72. E 73. A 74. C 75. E 76. E 77. C 78. D 79. B 80. A 81. C 82. B 83. D 84. E 85. B

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SAT CHEMISTRY Dr. D. Bampilis

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