i. chapter 9 the mole and stoich packet by ChemistryAdventure

The Mole

How heavy are molecules? Unit 9 In our last unit we learned how to mix chemicals and predict the products of the chemical reaction. What we didn’t learn is the amounts that are involved. That’s what this unit is all about. To do this will explore the mole, which is a unit of measure that enables us to find out exactly the amounts and the masses involved in any chemical reaction. Tentative Schedule Day 1: What is the mole and why is it important? Lab: Introduction to the mole Lesson: The Mole (slides 1-4) Homework: I’d like a mole worksheet 9.1 (in class) Reactions using the mole worksheet 9.2 Day 2: Using the mole Lab: Neutralization lab (slides 5-6) Lesson: Stoichiometry and the space shuttle (slides 7-8) Homework: Grams and moles worksheet 9.3 How to solve basic stoichiometry questions workseet 9.4 Day 3: Solving Stoichiometry Problems Lesson: Three types of stoichiometry problems (slides 9-14) Homework: Stoichiometry mixed problems worksheet 9.5 Day 4: Mixed Stoichiometry Problems Lesson: In class problem solving: More mixed stoichiometry problems worksheet 9.6 Tougher stoichiometry problems worksheet 9.7 Homework: How to solve limiting/excess reactant problems worksheet 9.8 Day 5: Yield and Percent Composition; Limiting Reactant Lesson: Yield and % Composition; Limiting Reactant (slides 15-19) Homework: Limiting reactant worksheet 9.9 Percent Composition worksheet 9.10 Day 6: Molar Volume and Review Lab: Molar Volume of Hydrogen Lesson: Review How to Ace it guides Homework: Review for mole and stoichiometry test Day 7: Test on the mole and stoichiometry

Name:_______________________________________

Period:______

Lab 9.1

Introduction to the mole lab. (10 Points) Goal: To learn the practical utility of the mole. Introduction: It would be very easy to mix chemicals together if all atoms had the same mass. For example if we needed to make some magnesium sulfide, where Mg + S  MgS

We could simply mass out 10 grams of magnesium and 10 grams of sulfur, cook them up, and we would have 20 grams of magnesium sulfide. However, in reality each atom of sulfur weighs about 25% more than each magnesium atom:

Our ten-grams-of-each recipe would have too many _______ atoms and not enough _________ atoms. The perfect ratio is 24.3 grams of magnesium for every 32.1 grams of sulfur. This mixture allows for every magnesium atom to react with the same number of sulfur atoms- none are wasted. We call this a stoichiometric ratio. It is simply based on the fact that some elements are heavier than others. 24.3 grams of magnesium is a mole of magnesium. 32.1 grams of sulfur is a mole of sulfur. These are handy numbers. Find them on the periodic table. They are simply the atomic masses of each atom expressed in grams, readily found on any periodic table. They each contain the same number of atoms: 6.02 x 1023 atoms. This is known as Avogadro’s Number. As you can see, the mole is a convenient way to count out huge numbers of molecules by weighing them on a balance. The same kind of relative mass problems occasionally happen in the real world- we need to combine tiny things – too small to count - in the right ratio by weighing them. Putting molecules together is in some ways like attaching a nut to a bolt. The parts have a structure that allows them to combine, and they do not have the same mass. This lab will show you these kinds of combinations can be performed on any scale.

Situation number one: Nuts and Bolts Puzzle Each of you will be given one nut and one bolt. You need to buy a lot of these: 10,000 nuts and 10,000 bolts. They combine in a 1:1 ratio: 1 nut + 1 bolt ď&#x192; 1 widget (ok, I made up the name) Think of a way to obtain 1000 nuts and 1000 bolts, not by counting them, but by weighing them. 1. Our plan (1 point): To create a pile of 10,000 nuts and 10,000 bolts without counting them, we will_______________________________________________________________________ __________________________________________________________________________ __________________________________________________________________________ __________________________________________________________________________ ________________________________________________________________________ 2. Our data (1 point): This should include the mass of one nut and one bolt.

3. Our calculations (1 point). Be sure a stranger could follow your logic.

4. Conclusion (1 point): Based on our data, to get 10,000 nuts and bolts we need to mass out ________ grams of nuts and _________ grams of bolts. Extrapolating 5 (1 point). If our boss needed 5 million nuts and 5 million bolts, no problem. We would simply mass out ______ grams of nuts and _______ grams of bolts. Situation number two: Paper clip puzzle. Large paper clips = oxygen Medium paper clips = carbon Small paper clips = hydrogen Assemble a molecule of ethanol, C2H6O (see drawing below) Assemble a molecule of acetic acid, C2H4O2 (See drawing below) Mass them. Mass ethanol:_______g Mass acetic acid__________g

These two molecules combine to form ethyl acetate: H H H C C O H

H O

H C C

O H

H H

ethanol

acetic acid

H C C H

ethyl acetate

H O H

water

Note the stoichiometric ratio: 1 molecule of ethanol combines with 1 molecule of ethyl acetate to form 1 molecule of ethyl acetate and 1 molecule of water. Using this data, we can combine any number of our paper clip molecules using mass: 6. (1 point). To make 10 molecules of ethyl acetate, combine _____ grams of ethanol with ______ grams of acetic acid 7. (2 points) Assemble the two products and show the masses of all four products below: H H H C C O H

H O

H C C

O H

H H

ethanol

acetic acid

_____ g

H C C

ethyl acetate _____ g

H O H

water _____ g

8. (1 points) Is the law of mass balance conserved? Please explain.

9. Please disassemble your molecules and return them to their cups, and have your lab stamped by your instructor. (1 point).

Lab9.2 Name__________________________Date_________________Period______________ Using the Mole: Neutralization Lab 5 Points Common acids and bases react to form a salt, and water. For example: HCl (aq) + NaOH (aq) ď&#x192; NaCl + H2O When equivalent amounts of acid and base react, the solution becomes neutral. This can easily be confirmed using pH paper. If we use too much acid or base the excess remains, and the solution is not neutralized. Again, we can find out with pH paper by checking the pH of the solution. We can use the mole to perform any reaction using exactly the right amounts. For example, in the above reaction, if we were given 36.5 grams of HCl (1 mole) we would need to use 40 grams (1 mole) of NaOH to neutralize it. Too much NaOH, and the solution is basic- it will turn pH paper blue. If we use too little, the solution remains acidic, and the pH paper will be red. In this experiment I will give each of you some acid (acetic acid), and you must use a base (baking soda) to neutralize it. Once you have done so, come up to me and I will test the pH of the solution. Your grade will be determined by the pH of your solution. Here is the chemical reaction:

H O

H O NaHCO3

H C C

O H

H baking soda _____ g _____ mol

acetic acid _____ g _____ mol

H C C

H O H

H ethyl acetate

water

CO2 carbon dioxide

_____ g

_____ mol

If time permits, we can use different bases to neutralize the solution. Instructions: Work in groups of 2. 1. Put on safety gear- goggles, gloves, and aprons. 2. I will pass out random amounts of baking soda (NaHCO3) to each group. 4. Complete the table above, and use it to Neutralize the solution using the correct amount of acetic acid. 5. Come up to me when you are ready to have the pH checked. 6. Clean up when you are finished

Name: ______________________________________

Date: ______

Period: _____

WS9.1

Iâ&#x20AC;&#x2122;d like a mole, please Chemists use the mole all the time to figure out how many atoms or molecules are in a substance. The mole is an amount (6.02 x 1023), also know as Avogadroâ&#x20AC;&#x2122;s number. More importantly, the average atomic mass in grams for each element has exactly one mole of atoms. For example, a mole of carbon has a mass of 12.01 g, and those 12.01 grams contain 6.02 x 1023 atoms of carbon. We can say that the molar mass of carbon is 12.011 grams per mole. If I needed half a mole of carbon I would mass out 6 grams of carbon, and that would contain 3.01 x 1023 atoms of C. Similarly, 24 g of C is 2 moles of carbon, and contains 1.2 x 1024 carbon atoms. Example: What is the molar mass of water, and how many water molecules would it contain? Solution: Water is H2O. We can find the molar mass by adding up the average atomic masses: 2(1.01) + 16.00 = 18.02 grams/mol. Since this is one mole of water, it will contain 6.02 x 1023 water molecules. ______________________________________________________________________________ Use these relationships to answer the questions below. 1. A mole of aluminum contains _____ g of aluminum, and also consists of ____________atoms of aluminum. If I needed 2 moles of aluminum, I would mass out ______ grams of aluminum, recognizing that it contains______________ aluminum atoms. 2. A mole of TNT (trinitrotoluene, C7H5N3O6) NO2 H

H O2N

NO2 H

has a mass of ________grams. If I had 0.1 moles of TNT it would have a mass of _________ grams, and would consist of _____________ molecules of TNT. H

3. The lightest element is ________;a molecule of this element has a molar mass of _________g/mol (remember, in this class, like in the real world, we will use the terms molar mass and mole interchangeably). 4. Determine the molar mass of each of the following molecules: A. Sugar (C6H12O6): B. The smell of fish: triethylamine:

N C. Carbon dioxide D. Water

Name_________________________ Period______________

WS9.2

Reactions using the mole The mole can be used to find out the how much material is required for a reaction, and how much it produces. For example, lets make a mole of water from the elements: 2 H2 (g) + O2  2 H2O (l) This equation tells us that 2 moles of molecular hydrogen react with one mole of molecular oxygen to form 2 moles of liquid water. But if we were to really to make water this way, how much of each reactant should we mix? We can look up the average atomic masses (the “red numbers” on our periodic table) to see the relative amounts needed and produced: ____ grams of H2 react with ____ g of O2 to form ____ grams of water. Since this reaction makes two moles of water, we have performed this reaction on a 2 mole scale. _____________________________________________________________________________ Show the amounts involved when each of these reactions is performed on a one mole scale; the first is done for you. 1.

Mn + O2  MnO2 55 g 32 g 87g

2. Ba + F2  BaF2 ___ ___ ____ 3. Mg + Se  MgSe __ ___ ____ Here are some that don’t react in a 1:1 ratio: 4. 2 H2 + O2  2 H2O (done for you) 4g 32g 36g 5. 2 Li + Cl2  2 LiCl ____ ___ ____ 6. 2Al + 3 Cl2  2 AlCl3 ____ ____ ____ 7. If I burn a mole of methane, how much carbon dioxide will it produce? CH4 + 2O2  CO2 + 2H2O 16g ____ ____ ____ 8. We can make chalk by reacting lime (calcium chloride) with soda ash (sodium carbonate). Fill in the amounts necessary to make one mole of chalk.

CaCl2 + Na2CO3  2 NaCl + CaCO3 ____ _____ ____ _____ 9. Try one based only on words: How would you make one mole of methane from the elements?

10. A perfume chemist needs 102 grams (1 mole) of ethyl propionate, which has an odor similar to bananas. How much propionic acid , and how much ethanol should he mix? CH3CH2CO2H + CH3CH2OH  Propionic acid ethanol _____________ _________ CH3CH2CO2CH2CH3 + H2O Ethyl propionate water ____________ ____

Name_________________________ Period_____________

How to Solve “Basic” Stoichiometry Questions

WS9.4

All of the following questions are based on lighting butane on fire, for which we can write the following balanced equation, :

2 C4H10 + 13 O2  8 CO2 + 10 H2O Note the molar ratio of 2:13:8:10. In the world of Stoichiometry we consider all of these to be “equivalent”: 2 moles C4H10 = 13 moles O2 = 8 moles CO2 = 10 moles H2O. Note how each example uses these conversion factors.

Type 1: Moles to Moles # of Steps: 1 1) Take your initial number of moles and use the balanced chemical equation to convert moles of the required chemical. Example : How many moles of O2 are required to produce 8 moles of H2O? (For all of these problems we will assume there is an excess of the other reactant that is needed, which in this case is H2) 8 moles H2O x 13 moles O2 = 10.4 moles O2 10 moles H2O

You try some: 1. How many moles of O2 are required to produce 240 moles of H2O?

2. How many moles of butane (C4H10) are needed to produce 25 moles of water?

Type 2: Moles to Grams and Grams to Moles Problems # of Steps: 2 These are like mole to mole problems, with the added step of adding up the molar masses of the substance in question to determine the number of grams needed or created. Example: How many grams of O2 are required to produce 9 moles of CO2? 9 moles CO2 x 13 moles O2 x 32 grams O2 = 468 grams O2 8 moles CO2 1 mole O2 You try one: How many grams of CO2 will be produced from 17 moles of O2?

Here is a grams to moles problem. They are just like moles to grams problems, but in a different order. Example: How many moles of C4H10 are required to produce 54 grams of H2O? 54 grams H2O x 1 mole H2O 18 grams H2O

2 moles C4H10 10 moles H2O

= 0.6 moles C4H10

You try one: How many moles of C4H10 are required to produce 100 grams of CO2?

Type 4: Grams to Grams # of Steps: 3 These are three- step problems. We have to convert from the grams given to moles, then from the moles of that substance to the equivalent number of moles of the desired substance, and then finally we convert from the moles of that substance to the equivalent number of grams.

To put it another way, we go from grams to moles, then moles to moles, then moles to grams.

Example: How many grams of H2O are produced from 40 grams of O2? 40 grams O2 x

1 mole O2 x 10 moles H2O x 18 grams H2O = 17.31 grams H2O 32 grams O2 13 moles O2 1 moles H2O

You try one: How many grams of butane (C4H10) are required to produce 25 grams of CO2? Type 5: Converting to molecules or atoms Since the mole is an amount,

Name_________________________ Period_____________

WS9.5

Stoichiometry Mixed Problems Example: For the reaction 2 H2 + O2  2 H2O, if you burn 10 grams of hydrogen with excess oxygen, how many moles of water will you make? Solution: this is a grams to moles problem :

10 grams H 2 x

mole H 2 2 moles H 2 O x  5 moles H 2 O 2 grams H 2 2 moles H 2

1. The reaction of magnesium sulfate with table salt produces magnesium chloride and sodium sulfate. Balance the reaction below: _____MgSO4 + _____NaCl  ____MgCl2 + ____Na2SO4

Use the equation in question 1 to solve, questions 2 and 3. 2. How many moles of NaCl must be used in order to produce 42.1 moles of Na2SO4?

3. How many moles of MgSO4 must be used in order to produce 100 moles of MgCl2?

4. _____V + _____O2  _____V2O5

Use the equation in question 4 to solve, questions 5 and 6. 5. How many moles of Vanadium are required to produce 47 grams of V2O5? Period: _____

6. How many grams of Oxygen gas are required to produce 31.4 grams of V2O5?

How to Solve Limiting/Excess Reactant Problems

WS9.8

In nature and in the lab, substance are being mixed together in all kinds of amounts. How much product will form? Here is a typical example. Example: If I mix together 10 grams of hydrogen with 10 grams of nitrogen, how many grams of ammonia (NH3) will I make? First we need to write a balanced chemical reaction: 3 H2 + N2  2 NH3 Next, we see how much product each reactant can make. The one that makes less is the limiting reactant. It’s like it sounds- it limits how much product can be made.

mole N2 2 mole NH3 17 g NH3 x x  12.1 g NH3 28 g N2 mole N2 mole NH3 mole H2 2 mole NH3 17 g NH3 10 g H2 x x x  56.7 g NH3 2 g H2 3 mole H2 mole NH3 10 g N2 x

This tells us that nitrogen is the limiting reactant. We can see how much H2 was wasted by seeing how much is needed:

10 g N2 x

2 g H2 mole N2 3 mole H2 x x  2.1 g H2 28 g N2 mole N2 mole H2

The rest of the hydrogen (7.9 g) is wasted. In summary, to perform limiting reactant problems 1. Write a balanced chemical reaction 2. Identify the limiting reactant by calculating how much product each reactant can make. The one that makes less is the limiting reactant. 3. To figure out how much excess reactant there is, see how much is needed and compare that to what you used. Use this example to guide you through these 2 problems. 1. Determine the mass of Tetraphosphorus Decoxide (P4O10) which is formed from 25.0 grams of Phosphorus (P4) and 50.0 grams of Oxygen gas (O2).

2. You start with 200.0 grams of Vanadium (V) and 100.0 grams of Oxygen Gas (O2) and the only product is Vanadium Oxide (V2O5). a) What is the limiting reactant? b) How many grams of Vanadium Oxide (V2O5) can form? c) How much excess reactant remains?

Limiting Reagent Worksheet

WS9.9

All of the questions on this worksheet involve the following reaction: When copper (II) chloride reacts with sodium nitrate, copper (II) nitrate and sodium chloride are formed. 1)

Write the balanced equation for the reaction given above:

If 15 grams of copper (II) chloride react with 20 grams of sodium nitrate, how much sodium chloride can be formed?

(13 g) 3)

What is the limiting reagent for the reaction in #2? __________________

How much of the excess reagent is left over in this reaction?

(0.88 g) If 11.3 grams of sodium chloride are formed in the reaction described in problem #2, what is the percent yield of this reaction? (87%)

Name__________________________Period______________

WS9.10

Percent Composition Worksheet

1. What is the percent composition of table salt (NaCl)

2A. What is the percent composition of carbon dioxide?

2B.How many grams of oxygen are present in an 88.0 g sample of carbon dioxide (CO2)?

3. What is the percent composition of sucrose (C12H22O11)?

4. What is the percent composition of aluminum sulfate (Al2(SO4)3)?

5. Which of the following compounds contains the highest percentage of iron? (Hint: compare the percent compositions) FeS2 Fe2O3 FeCO3

6. Caffeine has a chemical formula of C8H10N4O2. a) What is the molar mass of caffeine?

b) What is the percent composition of caffeine?

How to ace the mole test L1 Chemistry

In the previous unit (chemical reactions) we learned qualitatively how chemicals react. In this our 9th unit we took a quantitative look at chemical reactions by studying The Mole. We took a practical approach, emphasizing how the mole (usually referred to as molar mass) is used everyday to perform chemical reactions. The key idea is that each element has a unique characteristic mass- the average atomic mass. The mole is simply that mass in grams. Since that number is scaled to the mass of the substance, it means it will always equal the same number of particles. This number is Avogadro’s number: 6.02 x 1023. Practicing chemists very rarely use Avogadro’s number, and very often use molar mass. This unit emphasizes molar mass: the practical application of the mole. Still not clear? Well, the examples are easier to follow. 12.011 grams of pure carbon (for example a 12.011 gram diamond) is a mole of carbon and would contain 6.02 x 1023 atoms of carbon. 16 grams of methane (CH4) would be approximately one mole of methane (12 grams for carbon, and one gram for each hydrogen) and would contain 6.02 x 1023 molecules of methane. With the mole it is easy to create recipes that allow chemicals to react in exactly the right amounts, with nothing wasted. For example, reacting 12 grams of carbon (1 mole) with 32 grams of oxygen (that’s one mole of O2) will produce exactly one mole (44 g) of carbon dioxide. Using the mole we can perform any chemical reaction using just the right amount of reactants, so that nothing is wasted. We can also predict the exact amounts of each product formed. Using the mole we can also calculate the percent composition of any substance, a great tool for chemical identification. Water, for example, is always 89% oxygen, and 11% hydrogen. We know this because each mole of water must consist of one mole of oxygen atoms (16 g) and 2 moles of hydrogen (2 g)…so water is 16/18 oxygen, and 2/18 hydrogen by mass. **********************************************************************************************

To ace the mole test you should emphasize the following topics: 1. Know what the mole is- both in terms of how many particles, and in terms of mass.

Ex: a mole of carbon dioxide has a mass of ____ g and contains ______molecules.

2. Be able to measure out one mole of any substance, or any fraction thereof.

Ex: To give me a mole of water you would mass out ____ g.

3. Be able to perform a chemical reaction on any scale using the mole. This is perhaps the most common use of this handy topic.

Ex. To make 2 moles of water from the elements, the balanced chemical equation is: _________ + __________ --> _______ so I should mix ___g of hydrogen and ___g of oxygen

4. For any chemical reaction, be able to determine the number of molecules that are involved.

Ex. The above reaction uses ____molecules of hydrogen and ____molecules of oxygen to make ____ molecules of water.

5. Be able to predict the percent composition of any substance.

Ex. Ammonia (NH3): ___%N, ___%H

6. We should be familiar with all types of mole conversions: mole-mole (1 step), mole-gram or gram-mole (2 steps), and mole-mole (3 steps).

Examples: The combustion of one mole of hydrogen with excess oxygen will produce ____ moles of water.

The combustion of 4 moles of hydrogen with excess oxygen will produce _____ grams of water.

The combustion of 4 grams of hydrogen with excess oxygen will produce ___ grams of water. 7. We should be familiar with the concept of a limiting reactant. This is indispensable whenever any amounts of reactants are combined. Unless the stoichiometric amounts are used, this will create a situation where one reactant is in excess, and the other is limiting. We used a non-intuitive but useful method for solving these problems: we find out how much product each reactant will produce, and the lower amount “wins” – it defines the limiting reactant and therefore indicates how much product will form.

For example, what will happen when one gram of hydrogen is combined with one gram of oxygen? 1 g H2 x ____________ x___________ x _________ = g H 2O 1 g O2 x ____________ x ___________ x ___________ = g H2O The limiting reactant is _____, and this reaction will produce ______ g H2O. We learned a quick method for determining how much excess reactant there is- the ratio for the grams of product formed shows how much reactant was consumed. For the reaction above, since oxygen is the limiting reactant, the amount of hydrogen left is:

1 g H2 – 1g H2 (__/___) = _____ g excess hydrogen.

8. Finally, we learned to calculate the yield of a reaction; this is the actual yield/the theoretical yield x 100.

For example, if this reaction above produced 0.1 grams of water, the yield would be ( ___g/____g) x 100 = ____% yield.

Since this unit has a lot of mathematical conversions that follow the same series of steps, it is useful to drill them. There are also a couple of vocabulary problems. Practice them tonight. 1. What is stoichiometry? 2. Where does the word come from? 3. How to balance chemical reactions- try mixing barium chloride with sodium fluoride. Put your molar masses below to help answer all questions. Do this carefully- it will be used in several subsequent questions.

Molar masses:

__________ + ___________ --> ___________ + __________ Barium chloride sodium fluoride barium fluoride sodium chloride _____ g/mol ______ g/mol _____g/mol ________ g/mol

4. Coefficients (the numbers in front, like 2 H2O) vs. Subscripts (the little numbers, like 2 H2O)

Example: In the equation above put circles around each subscript, and squares around each coefficient.

5. Mole ratios : what are they for the above reaction?

___________ moles barium chloride = _______ moles sodium fluoride = _________ moles _____________ = _________ moles _______________

6. Mole-mole conversions

Example: 2.1 moles of barium chloride will make _____ moles of sodium chloride

7. Mole-mass conversions

Example: ten moles of barium chloride will make _____ grams of sodium chloride

8. Mass- mass conversions

Example: 100 grams of barium chloride will make ____ grams of sodium chloride.

9. Limiting reactants

Example: 100 grams of each reactant above will make _____ grams of sodium chloride. The limiting reactant is ____________.

10. Excess reactants.

In the example above, ________ is the excess reactant, and _____ grams of it remains in the pot after the reaction is complete.

11. Theoretical yield (this is 100% yield).

In the example above the theoretical yield of sodium chloride is ___ g.

12. Actual yield.

Example: If 50 grams of sodium chloride was obtained, that is a ____% yield.