Unit 8 Stoichiometry - North Allegheny · SECTION 12.2 CHEMICAL CALCULATIONS ... Since stoichiometry calculations are exactly the same thing as unit ... Answer the following questions

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Unit 8

Stoichiometry (Chapter 12)

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Name ____________________________________________________ Period ___________________

STOICHIOMETRY SECTION 12.1 THE ARITHMETIC OF EQUATIONS (pages 353–358) This section explains how to calculate the amount of reactants required or product formed in a nonchemical process. It teaches you how to interpret chemical equations in terms of interacting moles, representative particles, masses, and gas volume at STP. Using Everyday Equations (pages 353–355) 1. How can you determine the quantities of reactants and products in a chemical reaction?

____________________________________________________________________________________

______________________________________________________________________________

2. Quantity usually means the ______________________ of a substance expressed in grams or

moles.

3. A bookcase is to be built from 3 shelves (Sh), 2 side boards (Sb), 1 top (T), 1base (B), and 4 legs

(L). Write a “balanced equation” for the construction of this bookcase.

_________________________________________________________________________________

Using Balanced Chemical Equations (page 354) 4. Is the following sentence true or false? Stoichiometry is the calculation of quantities in chemical

reactions. ______________________

5. Calculations using balanced equations are called __________________________.

Interpreting Chemical Equations (pages 356–357) 6. From what elements is ammonia produced? How is it used? ____________________________

____________________________________________________________________________________

______________________________________________________________________________

7. Circle the letter of the term that tells what kind of information you CANNOT get from a

chemical equation.

a. moles b. mass c. size of particles d. volume e. number of particles

You can use the balanced equation. 3Sh _ 2Sb _ T _ B _ 4L _ Sh3Sb2TBL4 Ammonia molecules are composed of nitrogen and hydrogen; it is used as a fertilizer.

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CHAPTER 12, Stoichiometry (continued) 8. The coefficients of a balanced chemical equation tell you the relative number of moles of

______________________ and ______________________ in a chemical reaction.

9. Why is the relative number of moles of reactants and products the most important information

that a balanced chemical equation provides? _______________________________

____________________________________________________________________________________

______________________________________________________________________________

Mass Conservation in Chemical Reactions (pages 357–358) 10. Is the following sentence true or false? A balanced chemical equation must obey the law of

conservation of mass. ______________________

11. Use Figure 12.3 on page 357. Complete the table about the reaction of nitrogen and hydrogen.

12. Circle the letter(s) of the items that are ALWAYS conserved in every chemical reaction.

a. volume of gases d. moles

b. mass e. molecules

c. formula units f. atoms

13. What reactant combines with oxygen to form sulfur dioxide? Where can this reactant be found

in nature? __________________________________________________________________

_________________________________________________________________________________

_________________________________________________________________________________

atoms N + 6 atoms H ---> atoms N & atoms H

1 molecule N2 + molecules H2 ---> molecules NH3

x (6.02x1023 + 3 x (6.02x1023 ---> x (6.02 x 1023 molecules N2) molecules of H2) molecules NH3)

1 mol N2 + mol H2 ---> 2 mol NH3

28 g N2 + 3 x g H2 ---> 2 x g NH3

g reactants ---> 34 g products

N2 (g) + 2 H2 (g) ---> 2 NH3 (g)

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SECTION 12.2 CHEMICAL CALCULATIONS (pages 359–366) This section shows you how to construct mole ratios from balanced chemical equations. It then teaches you how to calculate stoichiometric quantities from balanced chemical equations using units of moles, mass, representative particles, and volumes of gases at STP. Writing and Using Mole Ratios (pages 359–362) 1. What is essential for all calculations involving amounts of reactants and products?

_________________________________________________________________________________

2. Is the following sentence true or false? If you know the number of moles of one substance in a

reaction, you need more information than the balanced chemical equation to determine the

number of moles of all the other substances in the reaction. ______________________

3. The coefficients from a balanced chemical equation are used to write conversion factors called

_________________________________ .

4. What are mole ratios used for? ______________________________________________________

____________________________________________________________________________________

____________________________________________________________________________________

___________________________________________________________________________

5. The equation for the formation of potassium chloride is given by the equation

2K(s) + Cl2(g) -->2KCl(s)

Write the six possible mole ratios for this equation.

______________________________ ______________________________

______________________________ ______________________________

______________________________ ______________________________

6. Is the following sentence true or false? Laboratory balances are used to measure moles of

substances directly. ______________________

7. The amount of a substance is usually determined by measuring its mass in_______________ .

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CHAPTER 12, Stoichiometry (continued) 8. Is the following sentence true or false? If a sample is measured in grams, molar mass can be

used to convert the mass to moles. ______________________

9. Complete the flow chart to show the steps for the mass–mass conversion of any given mass of

G to any wanted mass of W. In the chemical equation, a moles of G react with b moles of W.

10. Use the diagram below. Describe the steps needed to solve a mass–mass stoichiometry

problem. ________________________________________________________________________

____________________________________________________________________________________

______________________________________________________________________________

Other Stoichiometric Calculations (pages 363–366) 11. Is the following sentence true or false? Stoichiometric calculations can be expanded to include

any unit of measurement that is related to the mole. ___________________________

12. List two or three types of problems that can be solved with stoichiometric calculations.

____________________________________________________________________________________

____________________________________________________________________________________

mass of G = mol G molar mass G

x

mol G b mole W = a mole G

x

mol W molar mass W = mass W

x

aG bW (given quantity) (wanted quantity)

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Fun with Stoichiometry

The best definition for stoichiometry is this: It's a way of figuring out how much stuff you're going to make in a chemical reaction, or how much stuff you'll need to make a chemical reaction do what you want. It's very similar to making peanut butter and jelly sandwiches, except the conversion factors are a little bit more challenging. How do we get started? The first thing you must always write is a balanced chemical equation; then you will use stoichiometry calculations. Since stoichiometry calculations are exactly the same thing as unit conversions, let's review how to use the factor label method to convert units. Example: Convert 180 days of school to seconds.

Now that you remember how to do unit conversions, let’s review several vocabulary terms. Reactant: A substance that reacts in the reaction. For a chemical equation, the formulas on the _______________ side of the arrow are reactants. Product: A substance that you produce in the reaction. For a chemical equation, the formulas on the ______________ side of the arrow are products. Molar Mass: The mass in grams of one mole of a substance. This information can be obtained from the _____________________ _____________________. Coefficient: The number in front of a formula in the equation. It indicates the relative number of ______________ of each reactant and product in the reaction. Label the components of the following reaction with the correct vocabulary terms.

3Mg + Cr2S3 2Cr + 3MgS

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Fun with Stoichiometry The Mole Ratio

Mole Ratio: The ratio used to compare the number of moles of two substances in a chemical reaction. The numbers that fit into this ratio are pulled directly out of the _____________________ _____________________ for the reaction. The numbers that make up the mole ratio are the COEFFICIENTS in the balanced equation! For example, in the reaction: H2 + Cl2 2HCl, there are 6 different mole ratios that can be written.

**The mole ratios indicate the relative number of moles of each substance that are needed for the reaction to go to

_________________________________.**

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Fun with Stoichiometry ~ Follow the Boxes

The idea is simple. Given the problem you’re trying to solve, figure out what box in the diagram is your starting point and which box is the ending point. Then, use the factor label method to convert stepwise from the starting box through each intermediate step until the designation box is reached.

NEVER SKIP A BOX!!

EXAMPLE: mole mole

1. How many moles of oxygen are needed to produce 3.0 moles of water?

Balanced equation: _____________________________________________________

box ______ box _______

2. How many moles of hydrogen are needed to react completely with 20 moles of oxygen?


box ______ box _______

1 2 3 4

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EXAMPLE: mole gram

3. How many moles of hydrogen gas are needed to react completely with 20. grams of

chlorine?


box ______ box _______

4. How many grams of hydrochloric acid can be produced from 5.00 moles of chlorine?


box ______ box _______

EXAMPLE: gram gram

5. Copper reacts with a sufficient amount of silver nitrate. How many grams of copper

are required to produce 23.5 grams of silver?


box ______ box _______

6. How many grams of silver nitrate are required to react completely with 25 grams of

copper?


box ______ box _______

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Reaction Stoichiometry #1

In order to double or triple a recipe, the amount of EVERY ingredient had to be changed to keep the INGREDIENT RATIO constant. The same reasoning must be applied to a chemical equation. In order to compare reactants and products, you must know the “ingredient ratio,” which in a balanced chemical equation is the MOLE RATIO derived from the coefficients. Follow these basic steps to perform a stoichiometric calculation based upon a chemical reaction.

a. Write a balanced chemical equation. b. Determine your starting point box and ending point box. c. Solve the problem.

1. How many moles of silver can be recovered from the reaction between excess copper and 13.6

moles of silver nitrate?

STARTING BOX _____________ ------> ENDING BOX _____________

2. How many moles of hydrogen gas will react with 15.1 g of chlorine gas to produce hydrogen

chloride?


3. How many grams of calcium oxide are produced from the decomposition of 10.0 moles of

calcium carbonate?


4. How many grams of carbon dioxide are produced from the combustion of 1500 g of propane,

C3H8?


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5. Aspirin is synthesized by reacting salicylic acid (C7H6O3) with acetic anhydride (C4H6O3). The reaction is

2 C7H6O3 + C4H6O3 ---> 2 C9H8O4 + H2O What mass of aspirin, C9H8O4, can be produced from 500. g of salicylic acid?


6. The approximate annual consumption of gasoline in the U.S. is 146 billion gallons, which

averages to about 4 gallons a day per household. Assume gasoline to be 100% octane, C8H18. If this 146 billion gallons of gasoline has a mass of 3.8 x 1014 grams, how many grams of carbon dioxide are produced? (hint: the burning of octane is a combustion reaction)


7. Alka-Seltzer uses the reaction of sodium bicarbonate (baking soda) with citric acid (H3C6H5O7)

in aqueous solution to produce the fizz. The reaction is

3 NaHCO3 + H3C6H5O7 ---> 3 H2O + 3 CO2 + Na3C6H5O7 What mass of carbon dioxide is produced from 5 g (approximately 1 teaspoon) of baking soda? STARTING BOX _____________ ------> ENDING BOX _____________ 8. Silver jewelry tarnishes when the silver reacts with the sulfur found in the air. How many grams

of tarnish (silver sulfide) is formed from 0.05 g of sulfur?


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RXn Stoichiometry Prequiz #1 Name ____________________________ DIRECTIONS: Answer the following questions. You must use the factor-label method and show all of your work. 1. Write the balanced equation for the single replacement reaction between calcium and cuprous oxide.

___________________________________________________________________________________

a. How many moles of copper can be produced from 0.25 moles of cuprous oxide?

2. Write the balanced equation for the decomposition of sodium hydroxide.

__________________________________________________________________________________

a. How many grams of water are produced from the decomposition of 62.5 grams of sodium hydroxide.

3. Write the balanced equation for the composition of ammonia from its elements. ___________________________________________________________________________________

a. Write the balanced equation for the composition of methane from its elements.

______________________________________________________________________________

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Unit 8 ~ Problem Set #1 Pg. 360 #11, 12: Mole-Mole Calculations 11. This equation shows the formation of aluminum oxide, which is found on the surface of aluminum

objects exposed to the air. 4Al (s) + 3O2 (g) 2Al2O3 (s)

a. Write the six mole ratios that can be derived from this equation.

b. How many moles of aluminum are needed to form 3.7 mol Al2O3? box _____ box _____ 12. According to the equation in problem 11:

a. How many moles of oxygen are required to react completely with 14.8 moles of Al? box _____ box _____

b. How many moles of Al2O3 are formed when 0.78 moles of O2 reacts with aluminum? box _____ box _____

Pg. 362 #13, 14: Mass-Mass Calculations 13. Acetylene gas (C2H2) is produced by adding water to calcium carbide (CaC2).

CaC2 (s) + 2H2O (l) C2H2 (g) + Ca(OH)2 (aq)

How many grams of acetylene are produced by adding water to 5.00 g CaC2? box _____ box _____

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14. Using the same equation from #13, determine how many moles of CaC2 are needed to react completely with 49.0 g H2O. box _____ box _____

Pg. 364 #15, 16 15. How many molecules of oxygen are produced by the decomposition of 6.54 g of

potassium chlorate (KClO3)? 2KClO3 (s) 2KCl (s) + 3O2 (g)

16. The last step in the production of nitric acid is the reaction of nitrogen dioxide with water.

3NO2 (g) + H2O (l) 2HNO3 (aq) + NO (g)

How many grams of nitrogen dioxide must react with water to produce 5.00 x 1022 molecules of nitrogen monoxide?

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The Reactant’s the Limit! (but which one??) In carrying out a reaction, chemists rarely use exact stoichiometric mixtures of

reactants. They usually use an excess of one reactant to ensure that the reaction

goes to completion, and that the more expensive reactant is entirely consumed.

STEPS FOR COMPLETING A LIMITING REAGENT PROBLEM:

1. Convert grams of first reactant grams of desired product

2. Convert grams of second reactant grams of SAME product

a. smaller answer = MAXIMUM amount of product that can be produced

b. LR = reactant that produces the smallest amount

c. ER = the other reactant

3. Convert grams of LR grams of ER to determine amount ER you NEED

4. Excess = HAVE (original amount given) – NEED (calculated in #3)

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Example: Disulfur dichloride is used to vulcanize rubber. It can be made by a composition reaction between sulfur and chlorine gas. If you begin with 29.8 g of sulfur and 64.7 g of chlorine, how many grams of disulfur dichloride can be produced? How much of the excess reagent is left over?

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Name: ___________________________________________________________________ Period _______

Limiting Reagent Practice

1. You drop a 4.00 g hunk of zinc metal into 6.00 g of HCl in solution. How much hydrogen gas do you expect to produce from this reaction? What is the limiting reagent? What is the excess reagent? How much excess do you expect to have left over?

g of hydrogen gas: _____________ LR: _____________ ER: _____________ Left over: _____________ 2. In an extremely exothermic reaction, powdered aluminum metal will react with powdered iodine to

form aluminum iodide. If you react 5.00 g of aluminum with 20.0 g of iodine, how much aluminum iodide to you expect to produce from this reaction? What is the limiting reagent? What is the excess reagent? How much excess do you expect to have left over?

g of aluminum iodide:______________ LR: _____________ ER: ____________ Left over: _____________

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3. Being the pyro that she is, Miss Uhernik loves detonating pumpkins. To accomplish her explosion, Miss Uhernik adds 10.0 g of hydrogen gas and 5.00 g of oxygen gas to a balloon placed inside the gourd. How much water should be produced from this mixture? What is the limiting reagent? What is the excess reagent? How much excess do you expect to have left over?

g of water: ______________ LR: ______________ ER: ______________ Left over: ______________ 4. After overdoing it at wing night, you are experiencing some intestinal distress. To calm your stomach,

you take Maalox to neutralize the excess acid in your tummy. If you take 2.00 g of Maalox (magnesium hydroxide) to neutralize the 4.00 g of excess acid (HCl) in your stomach, how much magnesium chloride do you expect to produce? What is the limiting reagent? What is the excess reagent? How much excess do you expect to have left over?

g of Mg(OH)2: ______________ LR: ______________ ER: ______________ Left over: ______________

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Percent Yield

All of the quantities we have been calculating so far represent the maximum yield of product according to the balanced equation. Many reactions, however, fail to give a 100% yield of product. The main reasons for this failure are as follows: 1. Reactions do not always go to completion or may be reversible. 2. Impure reactants and competing side reactions may cause other products to form. 3. Some product may be lost in handling and transferring from one reaction vessel to another. The theoretical yield of a reaction is the calculated amount of product that should be obtained. In order to determine this amount, we must have a balanced chemical equation and do a stoichiometry problem. The actual yield of a reaction is the amount of product that actually forms when the reaction is carried out in the laboratory. An actual yield is an experimental value. The percent yield is the ratio of the actual yield to the theoretical yield multiplied by 100. Both the theoretical and the actual yields must have the same units to obtain a percent. Calculating percent yield measures the efficiency of the reaction in changing the reactants to products. % Yield = actual yield x 100 theoretical yield Example: What is the percent yield if 3.74 g of copper is produced when 1.87 g of aluminum is reacted with an excess of copper II sulfate? Example: Carbon disulfide can be made from sulfur dioxide and coke (coke is nothing more than carbon, C). Carbon dioxide gas is also produced. If the percent yield of carbon disulfide is 86%, how much will be produced from 950. g of coke?

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Percent Yield

1. When 50.0 g of silicon dioxide is heated with an excess of carbon, 32.2 g of silicon carbide is produced. SiO2 (s) + 3C (s) SiC (s) + 2CO (g)

What is the percent yield of this reaction?

2. The pollutant sulfur dioxide can be removed from the emissions of an industrial plant by reaction with calcium carbonate and oxygen.

2CaCO3 (s) + 2SO2 (g) + O2 (g) 2 CaSO4 (s) + 2CO2 (g)

If this reaction proceeds with a 96.8 % yield, how many kilograms of calcium sulfate are formed when 5.24 kg of sulfur dioxide reacts with an excess of calcium carbonate and oxygen gas?

3. When 84.8 g of iron III oxide reacts with an excess of carbon monoxide, 57.8 g of iron is produced.

The other product is carbon dioxide. What is the percent yield of this reaction?

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RXN Stoichiometry Prequiz #2 Name ____________________________ DIRECTIONS: SHOW ALL YOUR WORK for the following problems; report all answers to 3 significant figures. 1. Write the balanced equation for the reaction between aluminum and phosphoric acid.

a. How many grams of hydrogen are produced from a mixture of 15.0 grams of aluminum and 24.0

grams of phosphoric acid?

b. How much excess reagent is left over? 2. Write the balanced equation for the single replacement reaction between sodium and iron III oxide.

a. If 6.50 grams of iron are actually produced from 0.078 moles of iron III oxide, what is the percent

yield of the reaction?

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Name ____________________________________________________ Period ___________________

STOICHIOMETRY - Vocabulary Review

Match the correct vocabulary term to each numbered statement. Write the letter of the correct term on the line. Each answer can only be used once.

a. mole i. theoretical yield b. stoichiometry j. limiting reagent c. mass-mass calculation k. mole ratio d. reactants L. actual yield e. excess reagent m. percent yield f. atoms n. molar mass g. coefficient o. anion h. cation

___________ 1. the starting materials in a chemical reaction

___________ 2. a conversion factor derived from the coefficients of a balanced chemical

equation interpreted in terms of moles

___________ 3. the maximum amount of product that could be formed in a reaction

___________ 4. the amount of a substance that contains 6.02 x 1023 representative particles

of that substance

___________ 5. the substance completely used up in a chemical reaction

___________ 6. the ratio of how much product is produced compared to how much is

expected, expressed as a percentage

___________ 7. the calculations of quantities in a chemical reaction

___________ 8. the actual amount of product in a chemical reaction

___________ 9. the substance left over after a reaction takes place

___________10. a stoichiometric computation in which the mass of products is determined

from the given mass of reactants

___________11. mass and _______ are always conserved in a chemical reaction

___________12. the mass in grams of one mole of a substance

___________13. indicates the relative number of moles of each reactant and product in a

chemical reaction

___________14. a negative ion

___________15. a positive ion

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Reaction Stoichiometry Review

1. In the reaction Cu + 2 AgNO3 ----> Cu(NO3)2 + 2 Ag, the mole ratio of silver nitrate to silver is _____. 2. Consider the following equation:

3 Cu (s) + 8 HNO3 (aq) ----> 3 Cu(NO3)2 + 2 NO (g) + 4 H2O (l)

a. If 12 atoms of copper are dropped into a beaker of nitric acid, how many molecules of NO are produced? ____________

b. How many moles of water are produced along with 4.2 moles of copper II nitrate? ____________

c. How many grams of copper are needed to react with 8.6 moles of nitric acid? ______________ 3. Which of these is a correct interpretation of this balanced equation?

i. S (s) + 3 O2 (g) ----> 2 SO3 (g) a. 2 atoms S + 3 molecules O2 ----> 2 molecules SO3 b. 2 mol S + 3 mol O2 ----> 2 mol SO3

c. 6 mol S + 9 mol O2 -----> 6 mol SO3 d. all of the above 4. Is there a relationship between the total moles of reactants and total moles of products in a balanced

chemical equation? ______ 5. The amount of product that should be obtained under perfect conditions is called the ______________

________________. 6. The amount of product produced experimentally in the lab is called the _______________

________________. 7. In a chemical reaction, the limiting reagent is the reactant that is _______________ ________________. 8. In a chemical equation, the ______________ _________________ determines how much product can be

formed. 9. What two quantities are conserved in a chemical reaction? _______________ ________________. 10. Every reaction stoichiometry problem requires the use of a _______________ ________________. 11. Consider the reaction: 5A + 2B -----> 3C

When 10 moles of A react with an excess of B, 4 moles of C are actually formed. What is the % yield of the reaction?

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12. What chemical quantity is represented by the coefficients in a balanced equation? ________________ 13. The equation for the complete combustion of methane is

CH4 (g) + 2 O2 (g) -----> CO2 (g) + 2 H2O(l)

To calculate the number of moles of CO2 produced by the reaction of 50.6 g of CH4 with O2, the first conversion factor to use is:

a. 1 mol CH4 b. 16.0 g CH4

16.0 g CH4 1 mol CH4

c. 2 mol O2 d. 44 g CO2

1 mol CO2 2 mol CO2

14. List the three reasons why many reactions fail to give 100% yield.

_________________________________________________________________________________________________________________________________________________

_________________________________________________________________________________________________________________________________________________

_________________________________________________________________________________________________________________________________________________

15. A chemist interested in the efficiency of a chemical reaction would calculate the ________________

__________________. 16. Consider the equation: 3A + 2B ----> C

a. If the reaction stops when A is completely used up, then A is the _________________ ____________________.

b. If 10 grams of A completely react with 6 grams of B, how much product is formed? _________________________

17. Write the equation for the decomposition of lithium chlorate. 18. Write the equation for the reaction between hydrochloric acid and and cadmium. 19. Write the equation for the reaction between solutions of copper I nitrate and potassium iodide. 20. Write the balanced equation for the complete combustion of C3H8.

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DIRECTIONS: Show all of your work. Remember to use proper sig figs and units! 21. Write the balanced equation for the decomposition of chromium III chlorate.

How many moles of gas are produced when 29.6 grams of chromium III chlorate decompose? 22. Copper reacts with sulfur to form copper I sulfide. What is the maximum number of grams of copper I

sulfide that can be produced when 80.0 g of copper reacts with 25.0 g of sulfur? How much, in grams, of the excess reactant is left over?

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23. Molten iron and carbon monoxide are produced in a blast furnace by the reaction of ferric oxide with carbon. If 25.0 kg of pure ferric oxide are used, how many grams of iron can be produced?

24. Zinc was reacted with a solution containing 400. g of copper I sulfate. The reaction was stopped after

one hour, and 151 g of copper was obtained. Calculate the percent yield of copper obtained.

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Reaction Stoichiometry Pretest 1. The amount of product obtainable under perfect conditions is called the ________________________. 2. When Paul O’Neum weighs his final product and records his result, he has just recorded his

__________________ _________________. 3. In any chemical reaction the quantities that are conserved are ____________________ and

______________________. 4. In a chemical reaction the reactant that is completely consumed is called the ___________________

__________________ . 5. Tess Tube produces 25 lb of Smiley cookies, but according to the ingredients, Tess should be able to

produce 30 pounds. What is her % yield? Questions 6-9 refer to the following equation. Sodium hypochlorite, the active ingredient in bleach, is produced by the following reaction:

2 NaOH + Cl2 -----> NaCl + NaClO + H2O 6. How many moles of sodium hypochlorite (NaClO) can be produced from 10.2 moles of sodium

hydroxide?

7. How many grams of chlorine gas are needed to produce 14.8 moles of sodium hypochlorite? 8. How many grams of base are needed to produce 900. grams of table salt? 9. In the three problems, 6-8, what step is present in every one of them? ____________ _____________ 10. In order for Professor Brinclhof to measure the efficiency of a reaction, he would calculate his

________________ _________________. 11. If 16 grams of hydrogen gas react completely with 128 grams of oxygen gas, how many grams of water

will be produced? _____________. 12. When sodium bicarbonate is mixed with hydrochloric acid, the gas produced is tested with a flaming

splint. What is the result of the test? ________________________________________________________ 13. Name the seven diatomic elements. _________________________________________________________

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14. What is the molar mass of oxygen gas? ___________________________ 15. What is the formula for ammonia? _____________________________ 16. What is the charge of sodium in a compound? ___________ Oxygen? ___________

Calcium? ___________ Silver? ___________

17. The products of the decomposition of Na3N are ___________ and ___________. 18. Lake Erie is the fourth largest of the Great Lakes, with a surface area of 25 700 km2. How many sig. figs.

are in this measurement? _______ 19. Describe the relationship between moles of reactants and moles of products in a chemical reaction.

___________________________________________________________________________________________

_______________________________________________________________________________________

20. The coefficients in a balanced chemical equation describe the number of _________________ or number

of _____________________ of reactants and products. 21. Hydrofluoric acid cannot be stored in glass bottles because the acid reacts with silica in glass to produce

hexafluorosilicic acid: SiO2 + 6 HF -----> H2SiF6

+ 2 H2O If 30.0 grams of silicon dioxide and 40.0 grams of hydrofluoric acid react,

a. determine how much hexafluorosilicic acid is produced, b. determine how much excess reactant is left over, c. determine the percent yield if the actual yield is 45.8 grams of hexafluorosilicic acid.

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22. Write the balanced equation for the decomposition of magnesium chlorate. 23. Write the balanced equation for the single replacement reaction between potassium and silver nitrate. 24. Write the balanced equation for the double replacement reaction between hydrochloric acid and

sodium bromide. 25. Write the balanced equation for the complete combustion of acetylene (C2H2).

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Reaction Stoichiometry

ACROSS 1. Chemist’s “Bible,” Periodic

________

3. A product of a reaction

between plumbic nitrate

and zinc

5. A product resulting from

the decomposition of a

base

6. A unit used to measure a

quantity conserved in a

chemical reaction

8. Number of moles of H2 in

excess when 2 moles of H2

react with one

mole of O2 to form water

10. A product formed by the

addition of water to a

nonmetallic oxide

12. The step necessary in all

Reaction Stoichiometry

problems

15. Mercury I ion

16. A metallic bond is formed

between

positive ions and a ____ of

electrons

17. Study of chemical reactions

21. Formula for a possible

product of the reaction

between copper and

chlorine gas.

22. SO4-2

: sulfuric acid::

TeO4-2

: _______ acid

25. negative ions

27. Symbol for an inner

transition element in the

lanthanide series

28. Formula for the soluble

ionic salt formed from the

double replacement

reaction between aqueous

solutions of barium

chloride and potassium

sulftate

29. Place this noble gas in a

box (hint) and it may glow

bright orange; also an old

boxy Dodge

30. Least reactive metal

31. “Heart of Chemistry”,

abbreviation

33. Number of moles of

oxygen needed to react

with 184 grams of sodium

to form sodium oxide

34. The ____ reactant in a

reaction regulates the

______, or amount of

product obtained (2 words)

37. SI unit of mass

39. Old name for inner

transition elements: rare

______ elements

41. Acronym for light

amplification by stimulated

emission of radiation

42. Suffix that indicates one

less oxygen than “-ate”

43. Number of moles of Mg

with which nine moles of

Cl2 will react to form MgCl2

44. Nonmetals ______

electrons to form ions

45. Number of grams of

hydrogen gas formed from

the reaction between 80

grams of calcium and

excess hydrochloric acid

46. Solid, liquid, and gas: the

three _____ of matter

47. An alkali metal

48. Chem is ______

49. The symbol for the

feminine version of helium?

DOWN 1. The amount of product

that should be obtained

2. A metallic hydroxide

4. Reactant that is left over

7. Building block of matter

8. Metal used to coat nails

9. Rubidium bromide

11. SnF and NaF in toothpaste

help prevent _________

13. Substance found on the left

side of an equation

14. Poisonous

18. Amount of matter

19. The amount of product

obtained

20. Bitter: base:: _____:acid

23. Excess reactant

24. An example of evidence

that a chemical change has

occurred

26. Particle of a covalent

compound

31. Used to convert grams to

moles

32. Common name for NaOH

35. Nomenclature deals with

this branch of chemistry

36. Mass/Volume =

37. A sharp cutting instrument

made from carbide alloys

38. Each step of a reaction

stoichiometry problem

40. A use for silver bromide

44. A product formed from the

decomposition of

germanium chloride

46. Symbol for the element to

the left of yttrium

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Determination of the mole ratio

of reactants and products

DISCUSSION: An important application of the mole concept is expressing mole relationships between substances in a chemical reaction. According to the Law of Conservation of Matter, the total mass of the products in a chemical reaction must be the same as the total mass of the reactants. Also, the number of atoms in an ordinary chemical reaction must be conserved. Because atoms rearrange themselves into different combinations during a reaction, the total number of moles or molecules of products does not have to equal the total number of moles or molecules of reactants. Moles are not conserved. In this experiment, you will react sodium bicarbonate with hydrochloric acid. The products are sodium chloride, carbon dioxide, and water. The experimental determination of the relative masses of sodium bicarbonate and sodium chloride will enable you to determine the relative number of moles of reactants and products. Using this ratio you can write a balanced chemical equation for the reaction. OBJECTIVES:

to verify experimentally mass relationships and mole relationships between reactants and products of a chemical reaction

write balanced chemical equations and label a reaction by its type

MATERIALS: Bunsen burner evaporating dish wire gauze 10 mL graduated cylinder ring stand balance sodium bicarbonate hydrochloric acid, 3 M PROCEDURE: 1. Clean and dry and evaporating dish. 2. Place the evaporating dish on en electronic

balance and measure its mass. Record. _____________

3. Add one level spoonful of sodium bicarbonate to the evaporating dish. Measure the mass of the sodium bicarbonate and evaporating dish. Record. ___________

4. Slowly add 10 mL of hydrochloric acid to the sodium bicarbonate in the evaporating dish.

5. Then carefully add more acid, drop by drop, until the bubbling stops.

6. Using a ring stand apparatus and wire gauze, gently heat the evaporating dish with a small flame (about one inch below the wire gauze) until a dry solid remains.

7. Cool the evaporating dish and contents. Measure their mass and record. __________

8. Reheat the evaporating dish and contents with the gentle flame for five minutes.

9. Cool the evaporating dish and contents. Measure their mass and record. __________

10. Allow the evaporating dish to cool and then clean with water.

CONCLUSION: On a separate sheet of paper, answer the following questions. Be organized, neat, and label all parts. Pay attention to significant figures. 1. Complete the calculations on the following

page. 2. On the back of the calculations page,

construct a data table of all the measurements taken in the lab. Use a ruler and remember units and significant figures.

3. Under the data table, construct a calculations table to include your calculated answers. Use a ruler and remember units and significant figures.

Pop, Pop Fizz, Fizz

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Name______________________________________________________________ Period _________________ 1. Calculate the mass of sodium bicarbonate used in the reaction. 2. Calculate the moles of sodium bicarbonate used in the reaction. 3. Calculate the mass of the sodium chloride produced in the reaction. 4. Calculate the moles of sodium chloride produced in the reaction. 5. Calculate the experimental mole ratio of sodium bicarbonate to sodium chloride. Record the answer as

a decimal value. In other words, divide the denominator into the numerator. (2.35 moles sodium bicarbonate/1.62 moles sodium chloride = 1.45)

6. Write the balanced chemical equation for the reaction. 7. Using the balanced chemical equation, determine the actual mole ratio of sodium bicarbonate to

sodium chloride. 8. Calculate the percentage error for your experimental mole ratio: | actual - experimental | x 100 actual

9. Why is it important to heat the contents of the evaporating dish twice? ____________________________

___________________________________________________________________________________________

10. Explain the importance of using a gentle flame to heat the contents of the evaporating dish. __________

___________________________________________________________________________________________

____________________________________________________________________________________

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Determination of the percent yield of a reaction

DISCUSSION: In this experiment, a known amount of silver nitrate and copper metal will undergo a single replacement reaction. The two products will be separated, collected and weighed. Keep in mind that you will be graded on how well you recover the products, so be sure to follow the procedure carefully. OBJECTIVES: 1. To determine the actual mass of product. 2. To determine the theoretical mass of product. 3. To calculate the % yield of a reaction and %

error. MATERIALS: copper wire silver nitrate distilled water large test tube ring stand iron ring funnel filter paper 400 mL beaker balance PROCEDURE: DAY 1 1. Cut a piece of copper wire approximately 25

cm long. 2. Coil the lower part of the wire, leaving enough

straight wire to stick out of a large test tube. 3. Mass the copper wire to the nearest 0.01 g and

record the mass. __________________________ Set aside.

4. Clean a large test tube and rinse it with distilled water. Stand the clean test tube in a small beaker and place it on the balance; tare the balance.

5. Carefully spoon solid silver nitrate into the test tube until the mass is approximately 2.0 grams. (a little more or a little less if ok!) Record the mass. _________________ (CAUTION: Silver nitrate will react with your skin or clothing, turning it black. If you should get any on yourself, immediately wash it off.)

6. Fill the test tube about half way with distilled water. Swirl until all the silver nitrate has dissolved. (NOTE: If the water becomes cloudy, the test tube was not clean and you must start over.)

7. Put the coiled copper wire into the test tube. Make sure you have as much wire as possible in the solution.

8. Add distilled water until the water level is almost to the top.

9. Write your name on a piece of masking tape and place it near the mouth of the test tube. Place the test tube in the test tube rack designated for your class.

DAY 2 1. Using a pencil, place your name on a piece of

filter paper. Then weigh the paper to the nearest 0.01 g and record. ________________________ Make sure that you use the same balance as you used yesterday.

2. Set up the filtering apparatus; fold the filter paper and place it in the funnel. Moisten the paper with a few drops of distilled water to hold it in place.

3. Remove all the silver from the copper wire. a. Shake the silver crystals off the copper

wire into the test tube. Remove the wire.

b. Use your wash bottle to rinse onto the filter paper any remaining crystals adhering to the copper.

c. A few remaining crystals may need to be scraped from the wire with a stirring rod, then rinsed with more distilled water.

4. Set the copper wire on a clean paper towel and let it dry.

5. Pour the contents of the test tube into the funnel. Collect the blue solution (the filtrate) in the beaker.

6. Use distilled water to rinse any silver that was left behind in the test tube. Pour this rinse water onto the filter paper and let it drain into the beaker. Continue this process until all the silver has been transferred.

7. Wash the filter paper several times with distilled water.

8. Remove the filter paper from the funnel. Place it on a tray on the center table to dry overnight.

9. Pour the filtrate into the sink; clean the beaker. 10. Weigh the copper wire and record.

___________________ Discard the copper wire into the trash can.

DAY 3 1. Determine the mass of the silver and filter

paper. Record. __________________________ 2. Place the silver in the jar located on the central

distribution table and throw the filter paper into the garbage can.

“Au Gee.” It’s Only Silver

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“Au Gee, It’s Only Silver” Prelab

Earl N. Meyer dissolves 6.00 grams of silver nitrate in distilled water. He then adds 6.63 grams of

pure copper to the test tube. After the reaction is completed, he separates the leftover copper

from the silver and finds that 5.45 grams of copper is left over. After the silver dried overnight,

he weighed the silver and filter paper and found the mass to be 4.53 grams. The filter paper

weighs 1.02 grams.

1. Write the balanced chemical equation for the above reaction.

_________________________________________________________________________

2. What is the excess reactant? _____________________________________

3. What is the limiting reagent? _____________________________________

Data Table

Mass of copper before the reaction

Mass of silver nitrate used

Mass of filter paper

Mass of copper after the reaction

Mass of filter paper and silver

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“Au Gee, It’s Only Silver” Prelab 1. Calculate the mass of silver that was actually produced in Earl’s experiment. _______________________ 2. Calculate the mass of silver that should be (theoretical yield) produced in the experiment. _______________________ 3. Calculate the percent yield of his reaction. % yield = actual x 100

theoretical

_______________________ 4. Calculate the mass of copper that was actually consumed in his experiment. _______________________ 5. Calculate the mass of copper that should have (theoretical yield) reacted. _______________________ 6. Calculate the percent error in Earl’s amount of copper. % error = |theoretical – actual| x 100

theoretical

_______________________

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CONCLUSION: On a separate sheet of paper, answer the following questions. Be organized, neat, and label all parts. Pay attention to significant figures. 1. Construct a data table of all measurements

taken in the lab. Use the Prelab as a guide. Remember to use a ruler and units.

2. Calculate the following: a. the mass of silver nitrate actually produced b. the theoretical yield of silver c. the % yield of silver d. the mass of copper actually consumed e. the theoretical mass of copper consumed

f. % error of copper 3. Construct a Calculations table to include the

answers from a - f above. Remember to use a ruler and units.

4. Write a paragraph explaining the main source of error and how they affect your results. Remember, human error is not an excuse for poor results.

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Recycling old aluminum into alum crystals

INTRODUCTION:

Recycling old aluminum is mandatory in many cities and towns in the United States. Several states require a deposit on canned and bottled soft drinks to promote the recycling effort. Most states offer a certain price per pound on scrap aluminum. In this experiment, we will contribute to the recycling effort by using scrap aluminum as a starting material for the synthesis of alum, an inorganic compound. To produce the alum, the scrap aluminum will first react with potassium hydroxide according to the

following equation: 2Al (s) + 2KOH (aq) + 6H2O ---> 2K[Al(OH)4] (aq) +3H2 (g) Then the aqueous K[Al(OH)4], potassium tetrahydroxoaluminate III, will be added to sulfuric acid to produce the alum:

K[Al(OH)4] (aq) + 2H2SO4 (aq) + 8H2O –> KAl(SO4)2•12 H2O (s)

Alum is the common name for potassium aluminum sulfate dodecahydrate, KAl(SO4)2 • 12 H2O. The ionic hydrate contains 12 water molecules, six surrounding each of the two metallic cations. Alum has a number of useful applications. Most grocery stores have the substance on the shelves as a pickling agent to help pickles retain their crispness. Pool owners keep it on hand to clear up cloudy water in swimming pools. The textile industry uses it as a

bonding agent in dyes, and people use it every time the apply antiperspirant to block the flow of sweat glands.

OBJECTIVES:

1. To produce alum crystals from scrap aluminum. 2. To calculate the theoretical yield and actual yield

of a chemical reaction. MATERIALS: aluminum pie pan 1.4 M potassium hydroxide 9.0 M sulfuric acid scissors 2 250-mL beakers hot plate Buchner funnel Buchner flask filter paper graduated cylinders stirring rod 600 or 800 mL beaker

PROCEDURE:

DAY 1 1. Cut approximately 1 gram of aluminum from an

aluminum pie pan. Weigh it and record the mass. ________________

2. Cut the aluminum into tiny pieces, making sure

you do not lose any of it since you have already weighed the limiting reactant of the reaction.

3. Place the aluminum pieces in a clean 250-mL beaker; add

4. 50 mL of 1.4 M KOH. Bubbles of hydrogen gas should begin to form.

5. Heat the beaker gently on a hot plate (800C) to

speed up the reaction. (DO NOT use a Bunsen burner--the hydrogen gas that is generated is flammable!)

6. Heat until all the aluminum has been consumed. (No more bubbles!)

7. When you are finished heating the solution, the volume should be reduced to around 30 mL. If the volume falls below this level while heating, add distilled water to replenish it.

8. Set up a filtering apparatus; place a piece of filter paper in a Buchner funnel. Moisten the filter paper with a few drops of distilled water to hold it in place. Use rubber tubing to connect a Buchner flask to the aspirator. Place a 600 or 800

mL beaker in the sink directly below the aspirator and slowly fill it with water. This will help cushion the force of the water and prevent excessive splashing. Turn on the water full blast to create a vacuum in the flask.

9. Then and only then, filter your hot solution by slowly pouring it onto the filter paper. The filtrate should be clear, with any dark residue left on the filter paper. If it is not clear, filter again.

10. Rinse the beaker twice with 5 mL portions of distilled water. Pour each of the rinses through

the filter residue. 11. Transfer the clear filtrate into a clean 250-mL

beaker. Rinse the filtering flask with 5 mL of distilled water; pour the rinse into the beaker.

12. Discard the filter paper into the trash can. 13. When the filtrate is cool, slowly and carefully,

while stirring, add 20 mL of 9.0 H2SO4 to the solution.

14. As you are adding the sulfuric acid, a white precipitate should form. However, when all the acid has been added, the precipitate should

redissolve. 15. Warm the solution gently on a hot plate to insure

that all the precipitate has redissolved. The volume should be around 50 mL. If it is more than 50 mL keep heating until the volume is reduced (turn the hot plate up all the way to boil off all of the unwanted liquid).

16. STOP!! Write your name on the beaker and allow it to sit overnight in the hood to precipitate the alum crystals.

Junk Bonds

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DAY 2 1. Obtain your beaker and decant (pour out) the

liquid into the sink. Be careful not to lose any of your crystals.

2. Pour your crystals onto a paper towel and pat dry.

3. Place a small beaker on the balance you used on day 1 and tare. Slowly and carefully transfer the crystals from the paper towel into the beaker. Record the mass of the crystals. ____________________

PRELAB QUESTIONS: 1. Calculate the molar mass of CuSO4 • 5 H2O.

Follow the same procedure when calculating the molar mass of alum.

2. Add these two equations together to eliminate D, a product in the first equation and a reactant in the second.

A + 2B + 3C ---> 2D + E

D + 2F + 3C ---> 4G ∙ 9H2O

Just like in algebra, in order to eliminate D, the second equation must be doubled.

CONCLUSION: Answer the following on a separate sheet of paper. 1. Use a ruler to construct a data table that contains

the measurements taken in the lab.

2. Write the overall balanced equation for the synthesis of alum. This can be obtained by adding together the two equations found in the introduction. Use the prelab as a guide.

3. Using the overall equation determined in #2, calculate the theoretical yield of alum. Use the prelab as a guide when determining the molar mass of alum.

4. Calculate the percent yield of alum. 5. Use a ruler to construct a calculations table.

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WHAT dO I need to know?? Unit 8: Stoichiometry PART I: multiple choice

definitions (percent yield, actual yield, theoretical yield, limiting reagent, excess reagent, mole ratio, molar mass, law of conservation of mass)

calculations percent yield mole mole conversion atoms molecule conversion moles gram conversion

PART II: True & False

o percent yield, molar mass, diatomic elements, most wanted list, predicting products, limiting reagent, actual vs. theoretical yield, significant figures)

part III: Written Section

1. decomposition of a metallic chlorate a. grams reactant moles product

2. single replacement a. grams reactant grams reactant

3. double replacement a. percent yield

4. complete combustion a. grams reactant grams product b. excess remaining

5. predict products and write balanced chemical equations: a. double replacement (must label precipitate or circle the

molecular compound) b. complete combustion c. decomposition of a metallic chlorate d. single replacement reaction

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Documents

Unit 8 Stoichiometry - North Allegheny · SECTION 12.2 CHEMICAL CALCULATIONS ... Since stoichiometry calculations are exactly the same thing as unit ... Answer the following questions