Chemistry Name __________________________

Hour _______

Chemistry

Approximate Timeline

Students are expected to keep up with class work when absent.

CHAPTER 9 – CHEMICAL QUANTITIES

Day Plans for the day Assignment(s) for the day

1 Begin Chapter 9

Discuss Chapter 8 test

9.1 – Using Chemical Equations

o Information Given by Chem. Eq.

Identities of reactants/products

Physical state of reactants/products

Relative number of each molecule

o Mole-Mole Relationships

Leading into stoichiometry

Assignment 9.0

Assignment 9.1

Read section(s) 9.2

2 9.2 – Using Chem. Eq. to Calc. Mass

o Mass Calculations – stoichiometry

Assignment 9.2a

3 9.2 – Using Chem. Eq. to Calc. Mass

o Mass Calculations – stoichiometry

Assignment 9.2b

4 Discuss Experiment 37 – Stoichiometry

Assignment 9.2c due tomorrow

Assignment 9.2c

5 Perform Experiment 37

Lab Report due in 2 days

Read section(s) 9.3

6 9.3 – Limiting Reactants & Percent Yield

o Concept of Limiting Reactants

o Calculations with limiting reactants

Assignment 9.3a

7 9.3 – Limiting Reactants & Percent Yield

o Theoretical Yield v Actual Yield

o Percent Yield calculations

Assignment 9.3b

8 Work on Assignment 9.3c Assignment 9.3c

9 Work on Assignment 9.4 Assignment 9.4

10 Grade & discuss Assignment 9.4

Review for Chapter 9 Test

11 Chapter 9 Test Read section(s) 10.1

Chemistry Name __________________________

Hour _______

Study Guides

Chapter 9 Quizzes

Quiz 9.1 Using Chemical Equations 1. What information is given by a balanced chemical equation?

2. What is meant by the term “mole ratio”?

3. Use the mole ratio to answer questions about a chemical reaction.

4. Be able to convert between grams and moles.

Quiz 9.2 Using Chemical Equations to Calculate Mass 5. Define the term “stoichiometry”.

6. What are the three steps that constitute the “overall plan of attack” when solving

stoichiometry problems?

7. Solve a stoichiometry problem.

Quiz 9.3 Limiting Reactants & Percent Yield 8. Define the term “limiting reactant”.

9. After solving a stoichiometry problem, how do you know which reactant is the limiting

reactant?

10. How can you tell whether you are working a limiting reactant problem or a regular

stoichiometry problem?

11. Define the following terms.

a. Theoretical yield

b. Percent yield

12. What is the basic equation used to calculate percent yield?

13. Be able to solve a limiting reactant problem.

14. Be able to solve a percent yield problem.

Study Guide

Chapter 9 Test

At the completion of chapter 9 you should…

1. Know the definitions of the following terms.

a. Mole Ratio

b. Stoichiometry

c. Limiting Reactant

d. Theoretical Yield

e. Percent Yield

2. Understand the information given in a balanced chemical equation.

3. Give three ways to interpret a balanced chemical equation.

4. Use a balanced chemical equation to determine relationships between reactants and

products.

5. Relate masses of reactants and products (stoichiometry problems).

6. Understand the concept of limiting reactants.

7. Recognize the limiting reactant in a chemical reaction.

8. Perform limiting reactant calculations.

9. Understand percent yield.

10. Perform percent yield calculations.

Chemistry Name __________________________

Hour _______

Assignment 9.0 – Vocabulary

Define each of the following terms.

1. Mole Ratio

2. Stoichiometry

3. Limiting Reactant

4. Theoretical Yield

5. Percent Yield

Chemistry Name __________________________

Hour _______

Assignment 9.1 – Using Chemical Equations

1) What part of a chemical equation allows us to determine how much product we can get

from a known amount of reactants?

2) Why is it important to balance a chemical equation before proceeding to answer

questions about it?

3) Give at least three different ways to interpret the coefficients in a chemical equation.

4) How many moles of oxygen gas are required to react with 3.6 moles of sulfur dioxide in

the following reaction?

2 SO2 (g) + O2 (g) 2 SO3 (g)

5) Butane (C4H10) reacts with oxygen in the air to form carbon dioxide and water. How

many moles of carbon dioxide can be formed from 1.25 moles of butane and excess

oxygen?

2 C4H10 (g) + 13 O2 (g) 8 CO2 (g) + 10 H2O (l)

Chemistry Name __________________________

Hour _______

Assignment 9.2a – Stoichiometry (1) (2 pages)

Solve each of the following problems. Show all work. No work = no credit.

1) If 20.0 g of magnesium react with excess hydrochloric acid, how many grams of

magnesium chloride are produced?

_____ Mg + _____ HCl ---> _____ MgCl2 + _____ H2

2) How many grams of chlorine gas must be reacted with excess sodium iodide to produce

10.0 g of sodium chloride?

_____ NaI + _____ Cl2 ---> _____ NaCl + _____ I2

3) How many grams of oxygen are produced in the decomposition of 5.00 g of potassium

chlorate?

_____ KClO3 ---> _____ KCl + _____ O2

4) What mass of copper is required to replace silver from 4.00 g of silver (I) nitrate

dissolved in water?

_____ Cu + _____ AgNO3 ---> _____ Cu(NO3) 2 + _____ Ag

5) If excess ammonium sulfate reacts with 20.0 g of calcium hydroxide, how many grams of

ammonia are produced?

_____ (NH4) 2SO4 + _____ Ca(OH)2 ---> _____ CaSO4 + _____ NH3 + _____ H2O

Chemistry Name __________________________

Hour _______

Assignment 9.2b – Stoichiometry (2) (2 pages)

Solve each of the following problems. Show all work. No work = no credit.

1) If excess sulfuric acid reacts with 30.0 g of sodium chloride, how many grams of

hydrochloric acid are produced?

_____ NaCl + _____ H2SO4 ---> _____ HCl + _____ Na2SO4

2) How much silver (I) phosphate is produced if 10.0 g of silver (I) acetate reacts with

excess sodium phosphate?

_____ AgC2H3O2 + _____ Na3PO4 ---> _____ Ag3PO4 + _____ NaC2H3O2

3) How many grams of sodium hydroxide are needed to completely neutralize 25.0 g of

sulfuric acid?

_____ NaOH + _____ H2SO4 ---> _____ Na2SO4 + _____ H2O

4) Molten iron and carbon monoxide are produced in a blast furnace by the reaction of iron

(III) oxide and coke (carbon). If 2.5 x 104 g of pure Fe2O3 is used, how many grams of

iron can be produced?

_____ Fe2O3 + _____ C ---> _____ Fe + _____ CO

5) Ammonium nitrate, NH4NO3, is an important fertilizer and is also used in the

manufacture of explosives and fireworks. It is produced by treating nitric acid, HNO3,

with ammonia gas, NH3. If 340 kg of ammonia are used with all the nitric acid necessary

for reaction, how much ammonium nitrate can be produced?

_____ HNO3 + _____ NH3 ---> _____ NH4NO3

Extra Credit

6) Coal gasification is a process that is carried out industrially in a series of steps. The net

reaction involves coal (carbon) reacting with water to form methane, CH4, and carbon

dioxide. How many kilograms of methane can be produced from 1.00 x 103 kg of coal?

Chemistry Name __________________________

Hour _______

Assignment 9.2c – Using Chemical Equations to Calculate Mass (2 pages)

1) Why do we need to convert mass to moles in stoichiometry problems?

2) Calculate the number of grams of carbon dioxide that can be absorbed by complete

reaction with a solution that contains 5.00 grams of sodium hydroxide.

2 NaOH (aq) + CO2 (g) Na2CO3 (aq) + H2O (l)

3) You react 10.0 grams of nitrogen gas with hydrogen gas according to the following

reaction.

N2 (g) + 3 H2 (g) 2 NH3 (g)

A) What mass of hydrogen gas is required to completely react with 10.0 grams of

nitrogen gas?

B) What mass of ammonia is produced from 10.0 grams of nitrogen gas and excess

hydrogen gas?

4) You react hydrogen gas with chlorine gas according to the following reaction. What mass

of HCl (g) can be produced from 2.5 x 103 g of hydrogen gas with excess chlorine gas?

H2 (g) + Cl2 (g) 2 HCl (g)

5) Given the chemical equation below, determine the following.

CO (g) + 2 H2 (g) CH3OH (l)

A) What mass of carbon monoxide is required to produce 6.0 kg of methanol?

B) What mass of hydrogen is required to produce 6.0 kg of methanol?

Chemistry Name __________________________

Hour _______

Assignment 9.3a – Limiting Reactants (2 pages)

1) Fe2O3 + 2 Al ----> Al2O3 + 2 Fe

A) How many grams of iron are produced when 251.1 g of iron (III) oxide react with

34.02 g of aluminum?

B) What is the limiting reactant?

C) Which reactant is in excess?

2) 2 ZnS + 3 O2 ----> 2 ZnO + 2 SO2

A) What mass of sulfur dioxide is produced when 567 g of zinc (II) sulfide react with

409.0 g of oxygen?

B) What is the limiting reactant?

C) Which reactant is in excess?

3) N2 + 2 O2 ----> 2 NO2

A) What mass of nitrogen dioxide results when 14.8 grams of nitrogen react with

18.4 grams of oxygen?

B) What is the limiting reactant?

C) Which reactant is in excess?

Chemistry Name __________________________

Hour _______

Assignment 9.3b – Percent Yield

1) In the reaction between calcium hydride and water, the theoretical yield of hydrogen gas

is 7.18 grams. In running this reaction the actual amount of hydrogen produced was 6.94

grams. What was the percentage yield of this reaction?

2) Bubba was doing an experiment. He calculated that the reaction should produce 6.72

grams of product, but, after performing the lab, found that only 5.48 grams were

produced. What is the percentage yield?

3) Solid titanium(IV) oxide can be prepared by reacting gaseous titanium(IV) chloride with

oxygen gas. A second product of this reaction is chlorine gas.

TiCl4 + O2 TiO2 + 2 Cl2

A) Suppose that 6.71 x 103 g of titanium(IV) chloride is reacted with 2.45 x 10

3 g of

oxygen. Calculate the maximum mass of titanium(IV) oxide that can form.

B) If the percent yield of TiO2 is 75%, what mass is actually formed?

Chemistry Name __________________________

Hour _______

Assignment 9.3c – Limiting Reactants & Percent Yield (2 pages)

1) Suppose that a lemonade recipe calls for 1 cup of sugar for every 6 lemons. You have 12

lemons and 3 cups of sugar.

A) Which ingredient is limiting?

B) Why?

2) How is a limiting reactant problem different from other stoichiometry problems? (What is

you clue that it is a limiting reactant problem?)

3) A) What is the difference between theoretical yield and the actual yield?

B) Why is there a difference?

4) You react natural gas (methane, CH4) with chlorine gas to produce carbon tetra chloride

and hydrogen gas according to the following reaction: CH4 + 2 Cl2 CCl4 + 2 H2

Consider reacting 25.0 grams of each reactant and answer the following.

A) What mass of carbon tetrachloride should be produced (assuming 100% yield)?

B) Which reactant is limiting?

C) If only 22.2 g of carbon tetrachloride is obtained, what is the percent yield for this

reaction?

Chemistry Name __________________________

Hour _______

Assignment 9.4 – Chapter Review (6 pages)

1) True or false? For the reaction represented by 2 H2O2 (aq) 2 H2O (l) + O2 (g): if 2.0

grams of hydrogen peroxide decomposes, then 2.0 grams of water and 1.0 grams of

oxygen gas will be produced.

2) Consider the balanced equation CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g).

A) What is the mole ratio that would enable you to calculate the number of moles of

oxygen needed to react exactly with a given number of moles of methane?

B) What mole ratio would you need to calculate how many moles of each product

form from a given number of moles of methane? (two answers)

3) For each of the following balanced chemical equations, calculate how many moles and

how many grams of each product would be produced by the complete conversion of 0.50

mole of the reactant in boldface. State clearly the mole ratio used in each conversion.

A) NH3 (g) + HCl (g) NH4Cl (s)

B) CH4 (g) + 4 S (s) CS2 (l) + 2 H2S (g)

C) PCl3 (l) + 3 H2O (l) H3PO3 + 3 HCl (g)

D) NaOH (s) + CO2 (g) NaHCO3 (s)

4) What quantity serves as the conversion factor between the mass of a sample and the

number of moles the sample contains?

5) Calculate the mass of water vapor produced when 56.0 grams of pure hydrogen gas burns

in air.

2 H2 (g) + O2 (g) 2 H2O (g)

6) What mass of carbon dioxide is produced when 5.00 grams of carbon is burned with

excess oxygen?

C (s) + O2 (g) CO2 (g)

Chemistry Name __________________________

Hour _______

7) What mass of carbon monoxide is produced when 5.00 grams of carbon is burned with

excess oxygen?

2 C (s) + O2 (g) 2 CO (g)

8) Small quantities of oxygen gas can be generated in the laboratory by the decomposition

of hydrogen peroxide. The UNBALANCED equation is:

_______ H2O2 (aq) _______ H2O (l) + _______ O2 (g)

Calculate the mass of oxygen produced when 10.00 grams of hydrogen peroxide

decomposes.

9) Ammonium nitrate has been used as a high explosive because it is unstable and

decomposes into several gaseous substances. The rapid expansion of the gaseous

substances produces the explosive force. Calculate the mass of each product if 1.25

grams of ammonium nitrate reacts.

_______ NH4NO3 (s) _______ N2 (g) + _______ O2 (g) + _______ H2O (g)

10) How does the maximum yield of products for a reaction depend on the limiting reactant?

11) Suppose that exactly 50.0 grams of each reactant is taken.

_______ NH3 (g) + _______ Na (s) _______ NaNH2 (s) + _______ H2 (g)

A) Calculate the mass of the product in boldface produced when the limiting reactant

is completely consumed.

B) Which reactant is the limiting reactant?

12) Suppose that exactly 50.0 grams of each reactant is taken.

_______ BaCl2 (aq) + _______ Na2SO4 (aq) _______ BaSO4 (s) + _______ NaCl (aq)

A) Calculate the mass of the product in boldface produced when the limiting reactant

is completely consumed.

B) Which reactant is the limiting reactant?

Chemistry Name __________________________

Hour _______

13) If steel wool (iron) is heated until it glows and is placed in a bottle containing pure

oxygen, the iron reacts spectacularly to produce iron (III) oxide.

_______ Fe (s) + _______ O2 (g) _______Fe2O3 (s)

If 1.25 grams of iron is placed in a bottle containing 0.240 grams of oxygen gas, what

mass of iron (III) oxide will be produced?

14) According to his prelaboratory theoretical yield calculations, a student’s experiment

should have produced 1.44 grams of magnesium oxide. When he weighed his product

after reaction, only 1.23 grams of magnesium oxide was present. What is the student’s

percent yield?

15) Although they were formerly called inert gases, at least the heavier elements of Group 8

do form relatively stable compounds. For example, xenon combines directly with

elemental fluorine at elevated temperatures in the presence of a nickel catalyst.

Xe (g) + 2 F2 (g) XeF4 (s)

A) What is the theoretical mass of xenon tetrafluoride that should form when 130.

grams of xenon are reacted with 100. grams of fluorine?

B) What is the percent yield if only 145 grams of XeF4 is actually isolated?