chapters 9–12 resources - destiny high school · chapters 9-12 resources ... 4 chemistry: matter...

Chapters 9–12 Resources

Copyright © by The McGraw-Hill Companies, Inc. All rights reserved. Permission is granted to reproduce the material contained herein on the condition that such materials be reproduced only for classroom use; be provided to students, teachers, and families without charge; and be used solely in conjunction with the Glencoe Chemistry: Matter and Change program. Any other reproduction, for sale or other use, is expressly prohibited.

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ISBN: 978-0-07-878762-1MHID: 0-07-878762-9

Printed in the United States of America.

1 2 3 4 5 6 7 8 9 10 045 11 10 09 08 07

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To the Teacher . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . iv

Chapters 9-12 Resources

Reproducible Student Pages

Student Lab Safety Form . . . . . . . . . . . . . . . . . . . . . . . . . . vi

Chapter 9

Chemical Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . 1

Chapter 10

The Mole . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 33

Chapter 11

Stoichiometry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 57

Chapter 12

States of Matter . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 85

Teacher Guide and Answers

Chapter 9 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 109

Chapter 10 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 114

Chapter 11 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 118

Chapter 12 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 123

Table ofContents

iii

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Lab Safety Form

vi

Name:

Date:

Lab type (circle one) : Launch Lab MiniLab ChemLab

Lab Title:

Read carefully the entire lab and then answer the following questions. Your teacher must initial this form before you begin the lab.

1. What is the purpose of the investigation?

2. Will you be working with a partner or on a team?

3. Is this a design-your-own procedure? Circle: Yes No

4. Describe the safety procedures and additional warnings that you must follow as you perform this investigation.

5. Are there any steps in the procedure or lab safety symbols that you do not understand? Explain.

Teacher Approval Initials

Date of Approval

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Chapter 9 Chemical ReactionsMiniLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2

ChemLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3

Teaching Transparency Masters and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . . . 6

Math Skills Transparency Masters and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . . . 16

Study Guide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 20

Chapter Assessment. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 26

STP Recording Sheet . . . . . . . . . . . . . . . . . . . . . . . . . . . . 32

Table ofContents

1

Reproducible Pages

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2 Chemistry: Matter and Change • Chapter 9 ChemLab and MiniLab Worksheets

mini LAB 9Observe a Precipitate-Forming Reaction

Applying Concepts When two clear, colorless solutions are mixed, a chemical reactionmight occur, resulting in the formation of a precipitate.

Materials 150-mL beakers (2); 100-mL graduated cylinder; stirring rod (2); spatula (2);weighing paper (2); NaOH; Epsom salts (MgSO4�7H2O); distilled water; balance

Procedure1. Read and complete the lab safety form.

2. Place 50 mL distilled water in a 150-mL beaker.

3. Measure about 4 g NaOH pellets on a balance. Add the NaOH pellets to the beakerone at a time. Mix with a stirring rod until each NaOH pellet dissolves before addingthe next pellet.

4. Measure about 6 g Epsom salts (MgSO4) and place it in another 150-mL beaker. Add 50 mL distilled water to the Epsom salts. Mix with another stirring rod until theEpsom salts dissolve.

5. Slowly pour the Epsom salts solution into the NaOH solution. Record your observations.

6. Stir the new solution. Record your observations.

7. Allow the precipitate to settle, then decant the liquid from the solid into a 100-mL graduated cylinder.

8. Dispose of the solid as instructed by your teacher.

Analysis

1. Write a balanced chemical equation for the reaction between the NaOH and MgSO4.Most sulfate compounds exist as ions in aqueous solutions.

2. Write the complete ionic equation for this reaction.

3. Determine which ions are spectator ions, then write the net ionic equation for thisreaction.

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ChemLab and MiniLab Worksheets Chemistry: Matter and Change • Chapter 9 3

CHEMLAB 9

Safety Precautions• Always wear safety goggles and a lab apron.• Use caution when using sharp and coarse equipment.

ProblemWhich is the most reactivemetal tested? Which is theleast reactive metal tested?Can this information beused to predict whetherreactions will occur?

Objectives• Observe chemical

reactions.• Sequence the activities of

some metals.• Predict if reactions will

occur between certain substances.

Materials1.0M Zn(NO3)21.0M Al(NO3)31.0M Cu(NO3)21.0M Mg(NO3)2pipettes (4)wire cuttersCu wire

Al wireMg ribbonZn metal strips (4)emery cloth or fine

sandpaper24-well microscale

reaction plate

Develop an Activity SeriesSome metals are more reactive than others. By comparing how

different metals react with the same ions in aqueous solutions,an activity series for the tested metals can be developed. The activityseries will reflect the relative reactivity of the tested metals. It can beused to predict whether reactions will occur.

Pre-Lab

1. Read the entire CHEMLAB.

2. Make notes about procedures and safety precautions to use in the laboratory.

3. Use the data table on the next page.

4. Form a hypothesis about what reactions willoccur.

5. What are the independent and dependent variables?

6. What gas is produced when magnesium andhydrochloric acid react? Write the chemical equation for the reaction.

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7. Why is it important to clean the magnesium ribbon? How might not polishing a piece of metal affect the reaction involving that metal?

Procedure

1. Read and complete the lab safety form.

2. Use a pipette to fill each of the four wells in column 1 of the reaction plate with 2 mL of 1.0MAl(NO3)3 solution.

3. Repeat the procedure in step 2 to fill the fourwells in column 2 with 2 mL of 1.0M Mg(NO3)2solution.

4. Repeat the procedure in step 2 to fill the fourwells in column 3 with 2 mL of 1.0M Zn(NO3)2solution.

5. Repeat the procedure in step 2 to fill the fourwells in column 4 with 2 mL of 1.0M Cu(NO3)2solution.

6. With the emery paper or sandpaper, polish 10 cmof aluminum wire until it is shiny. Use wire cut-ters to cut the aluminum wire into four 2.5-cmpieces. Place a piece of the aluminum wire ineach row A well that contains solution.

7. Repeat the procedure in step 6 using 10 cm ofmagnesium ribbon. Place a piece of the Mg rib-bon in each row B well that contains solution.

8. Use the emery paper or sandpaper to polish smallstrips of zinc metal. Place a piece of Zn metal ineach row C well that contains solution.

9. Repeat the procedure in step 6 using 10 cm ofcopper wire. Place a piece of Cu wire in each row D well that contains solution.

10.Observe what happens in each cell. After 5 minutes, record your observations in the data table.

Cleanup and Disposal

1. Dispose of all chemicals and solutions as directedby your teacher.

2. Clean your equipment and return it to its properplace.

3. Wash your hands thoroughly before you leave thelab.

CHEMLAB 9

Al(NO3)3 Mg(NO3)2 Zn(NO3)2 Cu(NO3)2

Al

Mg

Zn

Cu

Reactions Between Solutions and Metals

Analyze and Conclude

1. Observe and Infer In which wells of the reaction plate did chemical reactions occur?Which metal reacted with the most solutions? Which metal reacted with the fewestsolutions? Which metal is the most reactive?

2. Sequence The most–active metal reacted with the most solutions. The least–active metalreacted with the fewest solutions. Order the four metals from the most active to the leastactive.

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3. Apply Write a chemical equation for each single-replacementreaction that occurred on your reaction plate.

4. Real–World Chemistry Under what circumstances might it beimportant to know the activity tendencies of a series of elements?

5. Error Analysis How does your answer from Question 2 above comparewith the activity series in Figure 9.13? What could account for thedifferences?

Inquiry Extension

Design an Experiment Think of three “what if” questions about this investigation thatmight affect your results. Design an experiment to test one of them.

CHEMLAB 9

METALSLithiumRubidiumPotassiumCalciumSodiumMagnesiumAluminumManganeseZincIronNickelTinLeadCopperSilverPlatinumGold

HALOGENSFluorineChlorineBromineIodine

Mostactive

Mostactive

Leastactive

Leastactive

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6 Chemistry: Matter and Change • Chapter 9 Teaching Transparency Masters

MgCl2(aq) � 2Na(s) 0 2NaCl(aq) � Mg(s)

2KI(aq) � Pb(NO3)2(aq) 0 PbI2(s) � 2KNO3(aq)

Reactants

ProductsSubscript Coefficients

Yields StateIn a watersolution

Equation 1:

6HCl(aq) � 2Al(s) 0 2AlCl3(Aq) � 3H2(g)

Equation 2:

Parts of a Balanced Chemical EquationParts of a Balanced Chemical Equation

TEACHING TRANSPARENCY MASTER

Use with Chapter 9,Section 9.1

29

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Teaching Transparency Worksheets Chemistry: Matter and Change • Chapter 9 7

Examine the parts of the chemical equation at the top of the transparency. Use thisinformation to answer the following questions about Equation 1 and Equation 2.

1. Write Equation 1 as a sentence.

2. Write Equation 2 as a sentence.

3. What substances are reactants in

a. Equation 1? b. Equation 2?

4. What substances are products in


5. List the coefficients used in

a. Equation 1. b. Equation 2.

6. What substances are in aqueous solution in


7. What substance shown is a gas?

8. What is the state of PbI2 in Equation 1?

9. What state is not represented in either equation?

10. What do the subscripts tell you in the formulas for

a. AlCl3?

b. KNO3?

c. Pb(NO3)2?

Parts of a Balanced Chemical EquationParts of a Balanced Chemical Equation

TEACHING TRANSPARENCY WORKSHEET


29

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Balancing Chemical EquationsBalancing Chemical Equations



30

Steps for Balancing Equations

1. Write the skeleton equation for the reaction.

2. Count the atoms of each element in the reactants.

3. Count the atoms of each element in the products.

4. Change the coefficients to make the number of atoms of each element equal on each side of the equation.

5. Write the coefficients in the lowest possible ratio.

6. Check your work.

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1. Examine the following equation.

Mg(s) � Ag2S(s) 0 MgS(s) � Ag(s)

a. How many atoms of magnesium are on each side of the equation?

b. Which element does not have the same number of atoms on both sides of the equation?

c. Write the balanced equation for this reaction.

2. Follow the steps for balancing a chemical equation and write a response for each step for the reaction in which iron metal (Fe) burns in oxygen (O2) to form iron(III)oxide (Fe2O3).

Step 1:

Step 2:

Step 3:

Step 4:

Step 5:

Step 6:

3. For each of the following, use at least one of the rules for balancing equations to explainwhy the equation is not properly balanced. Then write a correctly balanced equation foreach reaction.

a. 2H2O(l) � 2CO2(g) 0 2H2CO3(aq)

b. MgNO32(aq) � 2K(s) 0 Mg(s) � 2KNO3(aq)

c. AlCl3(aq) � AgNO3(aq) 0 AgCl(s) � Al(NO3)3(aq)

Balancing Chemical EquationsBalancing Chemical Equations



30

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The Activity SeriesThe Activity Series



31

Activity Series of Metals

Metal Reactions

LithiumRubidiumPotassiumCalciumSodium

MagnesiumAluminumManganeseZincIronNickelTinLead

CopperSilverPlatinumGold

All replace the hydrogen inwater and acids.

Each replaces metals listedbelow it.

All replace the hydrogen inacids.

Each replaces metals listedbelow it.

All are mostly unreactiveas far as replacing othermetals in a compound.

Activity Series of Halogens

Halogen Reactions

FluorineChlorineBromineIodine

Each replaces halogens listed below it.

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1. For each of the following pairs of elements, underline the one that would replace theother element in a compound.

a. calcium, tin e. iron, copper

b. bromine, fluorine f. iodine, chlorine

c. aluminum, potassium g. silver, lead

d. zinc, sodium

2. For each of the following reactants, use the activity series to determine whether the reaction would take place or not. If no reaction takes place, write NR in the blank. If areaction does take place, write the formulas for the products of the reaction. (Hint: If anactive metal replaces the hydrogen in water, the hydroxide of the active metal forms.)

a. Li(s) � Fe(NO3)3(aq) 0

b. Au(s) � HCl(aq) 0

c. Cl2(g) � KBr(aq) 0

d. Cu(s) � Al(NO3)3(aq) 0

e. Ag(s) � HBr(aq) 0

f. Ni(s) � SnCl2(aq) 0

3. Magnesium metal can be used to remove tarnish from silver items. Silver tarnish is thecorrosion that occurs when silver metal reacts with substances in the environment, especially those containing sulfur. Why would magnesium remove tarnish from silver?

4. Use the activity series for metals to explain why copper metal is used in plumbing wherethe water might contain compounds of many different metals.

5. The last four metals in the activity series of metals are commonly referred to as the“coinage metals.” Why would these metals be chosen over more active metals for use incoins? Why do you think some more active metals, such as zinc or nickel, are sometimesused in coins?

The Activity SeriesThe Activity Series



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Summary of Reaction TypesSummary of Reaction Types



32

Predicting Products of Chemical Reactions

Class of reaction Reactants Probable products

Synthesis Two or more substances One compound

Combustion A metal and oxygen The oxide of the metal

A nonmetal and oxygen The oxide of the nonmetal

A compound and oxygen Two or more oxides

Decomposition One compound Two or more elementsand/or compounds

Single-replacement A metal and a compound A new compound andthe replaced metal

A nonmetal and a A new compound and compound the replaced nonmetal

Double-replacement Two compounds Two different compounds, one ofwhich is often a solid, water, or a gas

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1. For each set of reactants listed below, identify the type of reaction that the reactantsmight undergo. List as many reaction types as may apply. Assume that all the reactantsfor the reaction are listed.

a. a compound and an element

b. two compounds

c. one compound

2. For each set of reaction products listed below, identify the type of reaction that mighthave formed the products. List as many reaction types as apply. Assume that all the products for the reaction are listed.

a. a compound and an element

b. two compounds

c. one compound

3. Classify each of the following examples according to the type of reaction involved. Listas many reaction types as may apply.

a. A match burns.

b. The carbonic acid found in soft drinks breaks down into bubbles of carbon dioxideand water.

c. Phosphorous and oxygen react rapidly, forming diphosphorous pentoxide.

d. An iron nail is placed into a copper sulfate solution. Copper metal appears on the nail.

e. The acid in baking powder reacts with baking soda (NaHCO3), forming carbondioxide gas and other products.

f. Water and sulfur trioxide react to form sulfuric acid.

g. Copper wire is placed in a silver nitrate solution. The solution turns blue, which is thecolor of the copper ion, and solid silver forms on the wire.

Summary of Reaction TypesSummary of Reaction Types



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Types of EquationsTypes of Equations



33

Co

mp

lete

eq

uat

ion

:

2KO

H(a

q)

�H

2SO

4(aq

) 0

K2S

O4(

aq)

�2H

2O(l

)

Co

mp

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ion

ic e

qu

atio

n:

2K�

(aq

) �

2OH

�(a

q)

�2H

�(a

q)

�SO

42�(a

q) 0

2K�

(aq

) �

SO42�

(aq

) �

2H2O

(l)

Net

ion

ic e

qu

atio

n:

OH

�(a

q)

�H

�(a

q) 0

H2O

(l)

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1. Write the complete equation as a sentence.

2. What is a spectator ion?

3. What are the spectator ions in this reaction?

4. Compare and contrast each pair below.

a. complete equations, complete ionic equations

b. complete ionic equations, net ionic equations

5. For the reaction between aqueous silver nitrate and aqueous sodium chloride, write eachof the following. The products of the reaction are aqueous sodium nitrate and solid silverchloride.

a. complete equation

b. complete ionic equation

c. net ionic equation

6. What is the net ionic equation for the reaction between aqueous calcium hydroxide and nitric acid? The products of this reaction are aqueous calcium nitrate and water. How does this net ionic equation compare to the net ionic equation shown on the transparency?

Types of EquationsTypes of Equations



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16 Chemistry: Matter and Change • Chapter 9 Math Skills Transparency Masters

Order of OperationsOrder of Operations

MATH SKILLS TRANSPARENCY MASTER


11

Order of Operations

Example Problem: 2 � 4 � 7 � (4 � 1)/3 � ?

Clear parentheses.

2 � 4 � 7 � (4 � 1)/3 � 2 � 4 � 7 � 3/3

Do any multiplication or division, from left to right.

2 � 4 � 7 � 3/3 � 2 � 28 � 3/3 � 2 � 28 � 1

Do any addition or subtraction, from left to right.

2 � 28 � 1 � 30 � 1 � 29

Order of Operations as Applied to Chemical Formulas

Example Problem: How many of each type of atom are present in2Al2(SO4)3? How many total atoms are present?

Clear parentheses.

In the (SO4)3 part of the formula, there are three units ofSO4, each with one S atom and four O atoms. So, (SO4)3contains three O atoms and 12 O atoms.

Do any multiplication or division, from left to right.

The coefficient of 2 must be multiplied times the number ofeach type of atom in the formula. Thus, there are 2 � 2, or4, Al atoms; 2 � 2, or 6, S atoms; and 2 � 12, or 24, O atoms.

Do any addition or subtraction, from left to right.

Total number of atoms: 4 � 6 � 24 � 34

Step 1

Step 2

Step 3

Step 1

Step 2

Step 3

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Math Skills Transparency Worksheets Chemistry: Matter and Change • Chapter 9 17

1. Follow the correct order of operations to solve these problems.

a. 4 � 2(9 � 1)

b. 9 � 14 � (4 � 3)

c. 7(2 � 14/2) � 23

d. 12/4(2 � 1 � 4 � 6)

e. 19 � 42/6 � 2(3 � 4 � 3 � 2)

2. Compare the answers to the following problems. Why is it important that parentheses be used?

a. 16/4 � 4 � ? b. 16/(4 � 4) � ?

3. For each of the following formulas, place parentheses where you think they should be.Explain your answers. If no parentheses are needed, write NP.

a. CaOH2

b. Na3PO4

c. AlNO33

d. NH42SO4

e. NiCO3

4. For each of the following formulas, list the number of each type of atom present and findthe total number of atoms represented.

a. KBr

b. Na2O

c. Au(OH)3

d. NH4NO3

e. (NH4)3PO4

f. 3K2SO4

g. 10C6H12O6

h. 6Mg3(PO4)2

Order of OperationsOrder of Operations

MATH SKILLS TRANSPARENCY WORKSHEET


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Simplifying EquationsSimplifying Equations



12

In an equation, both sides are equal. So, whatever is done tochange the terms on one side of the equation must also bedone to the other side.

Mathematical Equations:Simplifications such as combining like terms or clearing parentheses can be done on one side onlybecause they do not change the quantity present on that side; it is just expressed differently.

Like quantities can be added to or subtracted from bothsides of the equation. The entire equation can be multi-plied or divided by a number.

Chemical Equations:Net ionic equations are a simplified form of a complete chemical equation.

First, write the complete chemical equation as a complete ionic equation.

Then, simplify by subtracting like (spectator) ions from both sides.

Example Problem: Write the net ionic equation for thefollowing complete chemical equation.

NaOH � HNO3 0 NaNO3 � H2O

Write the complete ionic equation.

Na� � OH� � H� � NO3� 0 Na� � NO3

� � H2O

Subtract spectator ions from both sides of the equation.

Na� � OH� � H� � NO3� 0 Na� � NO3

� � H2O

OH� � H� 0 H2O

▲▲

▲▲

Step 1

Step 2

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1. Examine the following problem. Apply the rules for working with mathematical equations to correctly sequence the steps used to solve the problem. Write the numbers 1 through 6 in the blanks to show the correct order of the steps.

Solve for x in the following equation: 3(2x � 4) � 6 � 3x � 18.

6x � 3x � 6 � 6 � 3x � 3x � 18 � 6

6x � 12 � 6 � 3x � 18

6x � 6 � 3x � 18

3x/3 � 12/3

3x � 12

x � 4

2. Write balanced chemical, complete ionic, and net ionic equations for each of these chemical reactions.

a. Bubbles of carbon dioxide are released when nitric acid (HNO3) is added to a sodiumcarbonate solution. Water and sodium nitrate also form.

b. Bubbles of hydrogen sulfide are released when hydrochloric acid (HCl) is added to asolution of ammonium sulfide. Aqueous ammonium chloride also forms.

c. Potassium phosphate and water form when phosphoric acid (H3PO4) and potassiumhydroxide react.

d. Solid calcium sulfate and aqueous magnesium nitrate form when solutions ofmagnesium sulfate and calcium nitrate are mixed.

Simplifying EquationsSimplifying Equations



12

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20 Chemistry: Matter and Change • Chapter 9 Study Guide

Chemical ReactionsChemical Reactions

Section 9.1 Reactions and EquationsIn your textbook, read about evidence of chemical reactions.

For each statement, write yes if evidence of a chemical reaction is present. Write no ifthere is no evidence of a chemical reaction.

1. A tomato smells rotten.

2. A drinking glass breaks into smaller pieces.

3. A piece of ice melts.

4. Drain cleaner is mixed with water and the solution becomes warm.

5. Candle wax burns.

6. Molten candle wax solidifies.

7. Green leaves turn yellow and red as the seasons change.

8. Baking powder produces a gas that makes a cake rise.

In your textbook, read about how to represent chemical reactions and how to balancechemical equations.

Use the terms below to complete the passage. Each term may be used once, more thanonce, or not at all.

The fuel for the space shuttle is hydrogen, which burns in oxygen to produce water vapor

and energy. In this chemical reaction, hydrogen is a(n) (9) , oxygen

is a(n) (10) , and water vapor is a(n) (11) . In

a chemical equation for this reaction, a(n) (12) is used to separate

hydrogen and oxygen from water vapor and energy. A(n) (13) is

used to separate the symbols for hydrogen and oxygen. A(n) (14)

symbol is used to tell the state of hydrogen in the reaction, a(n) (15)

symbol is used for the state of oxygen, and a(n) (16) symbol is used

for the state of water vapor.

STUDY GUIDECHAPTER 9

arrow plus sign (s) (l)

reactant product (g) (aq)

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Study Guide Chemistry: Matter and Change • Chapter 9 21

For each of the following chemical reactions, write a word equation, a skeleton equation,and a balanced chemical equation. Be sure to show the state of each reactant and prod-uct. If you need more help writing formulas or determining the state of a substance,refer to Chapters 7 and 8 and the periodic table on pages 178–179.

17. Solid mercury(II) oxide breaks down when heated, forming the elements mercury andoxygen.

18. Sodium metal reacts with water vapor in air to form solid sodium hydroxide and hydrogen.

19. In the first step of refining zinc metal from its zinc sulfide ore, the ore is heated in thepresence of oxygen. The products are solid zinc oxide and sulfur dioxide gas.

20. The next step of refining zinc involves heating the zinc oxide in the presence of carbon.This reaction produces zinc vapor and carbon monoxide gas.

21. Certain pollutants in the air react with water vapor to form acids. For example, sulfur trioxide reacts with water vapor to form sulfuric acid.

22. Solid calcium carbonate is commonly used in antacids because it reacts with thehydrochloric acid found in the stomach. The products of this reaction are aqueous calcium chloride, carbon dioxide, and water.


Section 9.1 continued

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Section 9.2 Classifying Chemical ReactionsIn your textbook, read about synthesis, combustion, decomposition, and replacementreactions.

Assume that Q, T, X, and Z are symbols for elements. Match each equation in Column Awith the reaction type it represents in Column B.

Column A Column B

1. Q � XZ 0 X � QZ a. decomposition

2. Q � Z 0 QZ b. double-replacement

3. QT 0 Q � T c. single-replacement

4. QT � XZ 0 QZ � XT d. synthesis

Answer the following questions.

5. Does the following equation represent a combustion reaction, a synthesis reaction, orboth? Explain your answer.

2C(s) �O2(g) 0 2CO2(g) � energy

6. Why is it sometimes incorrect to state that a compound is broken down into its component elements in a decomposition reaction?

7. When soap is added to hard water, solid soap scum forms. When water is added to baking powder, carbon dioxide bubbles form. When lemon juice is added to householdammonia solution, water is one of the products. Tell how you know a double-replace-ment reaction has occurred in each case.

8. Explain how you can use an activity series to determine whether a single-replacementreaction will occur.


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In your textbook, read about the activity series for metals and halogens.

Examine each of the following pairs of potential reactants. Use Figure 9.13 in your textbook to help you decide whether or not a reaction would occur. If a reaction occurs,write the balanced equation. If no reaction occurs, write NR.

9. calcium and water

10. magnesium and water

11. rubidium and lithium chloride

12. potassium and aluminum oxide

13. silver and calcium nitrate

14. fluorine and potassium iodide

15. magnesium bromide and chlorine

16. copper and iron(III) sulfate

Match each example of a chemical reaction in Column A to the type(s) listed in Column B. List all types from Column B that apply.

Column A Column B

17. Aluminum lawn furniture becomes coated with a layer ofaluminum oxide when it sits out in the air.

18. Chlorine gas is bubbled through a calcium bromidesolution. The solution turns brown, the color of bromine.

19. Lime is added to acid water in a lake. Water and a saltform.

20. Propane is a common household fuel. When burned, waterand carbon dioxide are produced.

21. Steel wool burns, forming an iron oxide.

22. When an electric current is passed through moltenpotassium bromide, potassium and bromine form.

23. When solutions of sodium iodide and lead nitrate arecombined, a yellow solid forms.



a. combustion

b. decomposition

c. double-replacement

d. single-replacement

e. synthesis

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Section 9.3 Reactions in Aqueous SolutionsIn your textbook, read about aqueous solutions, reactions that form precipitates, reactions that form water, and reactions that form gases.

Circle the letter of the choice that best completes the statement or answers the question.

1. A spoonful of sodium chloride is dissolved in a liter of water. What is sodium chloride inthis solution?

a. molecule b. precipitate c. solute d. solvent

2. In an aqueous solution, water is the

a. homogeneous part. b. precipitate. c. solute. d. solvent.

3. Compounds that produce hydrogen ions in aqueous solutions are

a. acids. b. aqueous. c. bases. d. ionic compounds.

4. What type of reaction occurs between ions present in aqueous solution?

a. decomposition b. double-replacement c. single-replacement d. synthesis

5. What type of ions are present in solution but are not actually involved in a chemical reaction?

a. complete b. net c. precipitate d. spectator

6. If hydrochloric acid and potassium hydroxide react, what is the product of the net ionicequation for the reaction?

a. hydrochloric acid b. hydrogen ions c. potassium chloride d. water

7. Which of the following gases is not commonly produced in a double-replacement reaction?

a. carbon dioxide b. hydrogen cyanide c. hydrogen sulfide d. sulfur dioxide

8. H�(aq) � Br�(aq) � K�(aq) � OH�(aq) 0 H2O(l) � Br�(aq) � K�(aq) is an exampleof what type of chemical equation?

a. complete ionic b. net ionic c. precipitation d. spectator


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Predict the products for each reaction in Column A. Write the formulas for these prod-ucts on the product side of each equation. In the space at the left, write the letter of thechoice from Column B that indicates what type of product is produced during the reac-tion shown in Column A. Write as many choices as apply. (Hints: Compounds of group 1 metals are never precipitates; H2S and CO2 are gases.)

Column A Column B

9. HBr(aq) � KOH(aq) 0

10. HNO3(aq) � Na2CO3(aq) 0

11. NaI(aq) � Pb(C2H3O2)2(aq) 0

12. CsOH(aq) � H2SO4(aq) 0

13. K2S(aq) � HCl(aq) 0

For each of the following reactions, write chemical, complete ionic, and net ionic equations.

14. Phosphoric acid (H3PO4) and lithium hydroxide react to form a salt and water.

15. When solutions of magnesium sulfate and calcium chloride are mixed, calcium sulfateprecipitates.

16. Bubbles are released when nitric acid (HNO3) is added to a potassium carbonate solution.

17. Bubbles are released when hydrobromic acid (HBr) is added to a solution of ammoniumsulfide. Aqueous ammonium bromide also forms.



a. gas

b. precipitate

c. water

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Student Recording Sheet

32 Chemistry: Matter and Change • Chapter 9

Name Date Class

Standardized Test PracticeMultiple Choice

Select the best answer from the choices given, and fill in the corresponding circle.

1. 3. 5. 7.

2. 4. 6. 8.

Short Answer

Answer each question with complete sentences.

9.

10.

11.

Extended Response


12.

13.

14.

SAT Subject Test: Chemistry

15. 17.

16. 18.

CHAPTER 9

Assessment

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Chapter 10 The MoleMiniLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 34

ChemLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 35


Math Skills Transparency Masters and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . . 40

Study Guide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 44



Table ofContents

33

Reproducible Pages

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mini LAB 10Analyze Chewing Gum

Interpreting Data Are sweetening and flavoring added as a coating or mixed through-out chewing gum?

Materials balance, weighing paper, 250-mL beakers (2), pieces of chewing gum (2), stirring rod, paper towels, window screen (10 cm � 10 cm), scissors, clock or timer

Procedure CAUTION: Do not taste or eat any items used in the lab.


2. Unwrap two pieces of chewing gum. Place each piece on a weighing paper.Measure and record each mass using a balance.WARNING: Do not eat any items used in the lab.

3. Add 150 mL of cold tap water to a 250-mL beaker. Place one piece of chewing gum in the water, and stir with a stirring rod for 2 min.

4. Pat the gum dry using paper towels. Measure and record the mass of thedried gum.

5. Use scissors to cut the second piece of gum into small pieces. Repeat Step 3using fresh water. Keep the pieces from clumping together.WARNING: Use caution with scissors.

6. Use a 10-cm x 10-cm piece of window screen to strain the water from thegum. Pat the gum dry using paper towels. Measure and record the mass ofthe dried gum.

Analysis

1. Calculate For the uncut piece of gum, calculate the mass of sweeteners and flavor-ings that dissolved in the water. The mass of sweeteners and flavorings is the differ-ence between the original mass of the gum and the mass of the dried gum.

2. Calculate For the gum that was in small pieces, calculate the mass of dissolved sweet-eners and flavorings.

3. Apply For each piece of gum, determine the percent of the original mass from thesoluble sweeteners and flavorings.

4. Infer What can you infer from the two percentages? Is the gum sugar-coated or arethe sweeteners and flavorings mixed throughout?

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CHEMLAB 10

Safety Precautions • Always wear safety goggles and a lab apron.• Hot objects will not appear to be hot.• Use the Bunsen burner carefully. • Turn off the Bunsen burner when not in use.

ProblemHow can you determinethe moles of water in amole of a hydrated compound?

Objectives• Heat a known mass of

hydrated compound untilthe water is removed.

• Calculate the formula for a hydrate using themass of the hydratedcompound and the mass of the anhydrouscompound.

MaterialsBunsen burnerring stand and ringcrucible and lidclay trianglecrucible tongsbalanceEpsom salts (hydrated MgSO4)spatulaspark lighter or matches

Hydrated CrystalsHydrates are compounds that incorporate water molecules in their

crystalline structures. The ratio of moles of water to one mole ofthe compound is a small whole number. For example, in the hydratedcompound copper(II) sulfate pentahydrate (CuSO4�5H2O), the ratio is 5:1. The ratio of moles of water to one mole of a hydrate can bedetermined experimentally by heating the hydrate to remove water.

Pre-Lab


2. Prepare all written materials that you will takeinto the laboratory. Be sure to include safety pre-cautions and procedure notes. Use the data tableon the next page.

3. Explain how you will obtain the mass of waterand the mass of anhydrous MgSO4 contained inthe hydrate.

4. How will you convert the masses of anhydrousMgSO4 and water to moles?

5. How can you obtain the formula for the hydratefrom the moles of anhydrous MgSO4 and themoles of water?

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Procedure

1. Measure to the nearest 0.01 g the mass of a clean,dry crucible with a lid. Record the mass.

2. Add about 3 g hydrated MgSO4 to the crucible.Measure the mass of the crucible, lid, and hydrateto the nearest 0.01 g and record the mass.

3. Record your observations of the hydrate.

4. Place the triangle on the ring of the ring stand.Carefully place the crucible in the triangle.

5. Place the crucible lid on the crucible slightlycocked to help prevent spattering and allow vaporto escape. Begin heating with a low flame, thengradually progress to a stronger flame. Heat forabout 10 minutes.

6. When heating is complete, remove the crucibleusing tongs. Place the lid on the crucible andallow the crucible and contents to cool.

7. Measure the mass of the crucible, lid, and MgSO4and record the mass in the data table.

8. Observe the anhydrous MgSO4 and record yourobservations.


1. Discard the anhydrous MgSO4 in a trash con-tainer or as directed by your teacher.

2. Return all lab equipment to its proper place andclean your lab station.

3. Wash your hands thoroughly when all lab workand cleanup are complete.

CHEMLAB 10

Observations of hydrated MgSO4

Mass of crucible and lid

Mass of crucible, lid, and hydrated MgSO4

Mass of hydrated MgSO4

Mass of crucible, lid, and anhydrous MgSO4

Mass of anhydrous MgSO4

Mass of water in hydrated MgSO4

Moles of anhydrous MgSO4

Moles of water in hydrated MgSO4

Observation of anhydrous MgSO4

Mass Data and Observations of Epsom Salts

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1. Calculate Use your experimental data to calculate the formula for hydrated MgSO4.

2. Observe and Infer How do appearances of the hydrated and anhydrous MgSO4 crystalscompare? How are they different?

3. Conclude Why might the method used not be suitable for determining the water ofhydration for all hydrates?

4. If the hydrate’s formula is MgSO4�7H2O, what is the percent error forhydrated MgSO4? What are the possible sources for the error? What procedural changescould you make to reduce the error?

5. Predict the result of leaving the anhydrous crystals uncovered overnight.

Inquiry Extension

Design an Experiment to test whether a compound is hydrated or anhydrous.

Error Analysis

CHEMLAB 10

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Mas

s o

f co

mp

ou

nd

Mo

les

of

com

po

un

d

Mo

les

of

ato

ms

or

ion

s

Rep

rese

nta

tive

par

ticl

es

6.02

� 1

023

mo

lecu

les

or

form

ula

un

its

6.02

� 1

023

mo

lecu

les

or

form

ula

un

its

nu

mb

er o

f g

ram

s1

mo

l

1 m

ol

nu

mb

er o

f g

ram

s1

mo

l

1 m

ol

mol atoms or ions1 mol compound

1 mol compoundmol atoms or ions

Mass-to-Mole andMole-to-ParticlesConversions for Compounds

Mass-to-Mole and Mole-to-Particles Conversions for Compounds



34

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1. According to the diagram, what three quantities can you calculate if you know the number of moles of a compound?

2. According to the diagram, what three quantities can you calculate from a mass measurement of a compound?

3. If you were given the number of moles of a compound, what quantity would you need toknow to calculate the mass of that number of moles of the compound?

4. If you were given the number of moles of a compound, what information would youneed to know to determine each of the conversion factors necessary to find the number ofmoles of each atom or ion in the compound?

5. You are given a 2.0-mol sample of calcium carbonate (CaCO3). The molar mass ofCaCO3 is 100.09 g/mol. Write the conversion factor you would use to determine correctly each of the following quantities.

a. the mass in grams of the sample

b. the number of formula units of CaCO3 in the sample

c. the number of moles of oxygen atoms in the sample

6. Write the conversion factors in the order you would use them to determine correctly eachof the following quantities in a sample of 2.0 � 1024 molecules of ethane (C2H6). Themolar mass of ethane is 30.08 g/mol.

a. the mass in grams of the sample

b. the number of carbon atoms in the sample

Mass-to-Mole andMole-to-ParticlesConversions for Compounds

Mass-to-Mole and Mole-to-Particles Conversions for Compounds



34

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Calculations Involving the Molar Mass of an ElementCalculations Involving the Molar Mass of an Element



13

Calculating the Mass of a Given Number of Moles of an Element

What mass of carbon (C) atoms contains 3.00 mol C?

3.00 mol C � � 3.00 mol C � C � 36.0 g C

Calculating the Number of Moles in a Given Mass of an Element

How many moles of potassium (K) atoms are contained in 50.0 g K?

50.0 g � � 50.0 g � � 1.28 mol K1 mol K��39.098 g

1 mol K��39.098 g

12.011 g C��

1 mol C12.011 g C��

1 mol C

6

CCarbon12.011

[He]2s22p2

19

KPotassium

39.098[Ar]4s1

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1. Determine the mass in grams of each of the following. Use the periodic table.

a. 1.00 mol silver (Ag)

b. 12.0 mol aluminum (Al)

c. 3.25 mol copper (Cu)

d. 1.93 mol xenon (Xe)

e. 5.34 mol vanadium (V)

2. Determine the number of moles in each of the following. Use the periodic table.

a. 10.0 g lithium (Li)

b. 367 g magnesium (Mg)

c. 72.1 g silicon (Si)

d. 4.87 g fluorine (F)

e. 1.56 kg lead (Pb)

Calculations Involving the Molar Mass of an ElementCalculations Involving the Molar Mass of an Element



13

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Calculating the Molar Mass of a CompoundCalculating the Molar Mass of a Compound



14

What is the molar mass of gallium hydride (Ga2H6)?

1 mol Ga2H6 � � 1 mol Ga2H6 � � 2 mol Ga

1 mol Ga2H6 � � 1 mol Ga2H6 � � 6 mol H

2 mol Ga � � 2 mol Ga � � 139.446 g Ga

6 mol H � � 6 mol H � � 6.048 g H

molar mass Ga2H6 � 145.494 g/mol Ga2H6

1.008 g H��1 mol H

1.008 g�1 mol H

69.723 g Ga��

1 mol Ga69.723 g��1 mol Ga

6 mol H��1 mol Ga2H6

6 mol H��1 mol Ga2H6

2 mol Ga��1 mol Ga2H6

2 mol Ga��1 mol Ga2H6

1

HHydrogen

1.0081s1

31

GaGallium69.723

[Ar]4s23d104p1

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Determine the molar mass of each of the following compounds. Use the periodic table.

1. carbon dioxide (CO2)

2. mercury(I) fluoride (Hg2F2)

3. magnesium thiotellurite (Mg3TeS5)

4. copper(II) cyanide (Cu(CN)2)

5. cobalt(II) orthophosphate (Co3(PO4)2)

Calculating the Molar Mass of a CompoundCalculating the Molar Mass of a Compound



14

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The MoleThe Mole

Section 10.1 Measuring MatterIn your textbook, read about counting particles.

In Column B, rank the quantities from Column A from smallest to largest.

Column A Column B

0.5 mol 1.

200 2.

5 3.

6,000,000,000 4.

6.02 � 1023 5.

dozen 6.

four moles 7.

gross 8.

pair 9.

ream 10.

In your textbook, read about converting moles to particles and particles to moles.

In the boxes provided, write the conversion factor that correctly completes each problem.

11. 1.20 mol Cu � � 7.22 � 1023 Cu atoms

12. 9.25 � 1022 molecules CH4 � � 1.54 � 10�1 mol CH4

13. 1.54 � 1026 atoms Xe � � 2.56 � 102 mol Xe

14. 3.01 mol F2 � � 1.81 � 1024 molecules F2


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Section 10.2 Mass and the MoleIn your textbook, read about the mass of a mole.

For each statement below, write true or false.

1. The isotope hydrogen-1 is the standard used for the relative scale ofatomic masses.

2. The mass of an atom of helium-4 is 4 amu.

3. The mass of a mole of hydrogen atoms is 1.00 � 1023 amu.

4. The mass in grams of one mole of any pure substance is called its molarmass.

5. The atomic masses recorded on the periodic table are weighted averagesof the masses of all the naturally occurring isotopes of each element.

6. The molar mass of any element is numerically equal to its atomic mass in grams.

7. The molar mass unit is mol/g.

8. If the measured mass of an element is numerically equal to its molar mass,then you have indirectly counted 6.02 � 1023 atoms of the element in themeasurement.

In your textbook, read about using molar mass.

For each problem listed in Column A, select from Column B the letter of the conversionfactor that is needed to solve the problem. You may need to use more than one conver-sion factor to solve the problem.

Column A Column B

9. Find the number of moles in 23.0 g of zinc.

10. Find the mass of 5.0 � 1020 zinc atoms.

11. Find the mass of 2.00 moles of zinc.

12. Find the number of atoms in 7.40 g of zinc.

13. Find the number of moles that contain 4.25 � 1027 zincatoms.

14. Find the number of atoms in 3.25 moles of zinc.


a.

b.

c.

d.1 mol Zn

��6.02 � 1023 atoms Zn

6.02 � 1023 atoms Zn��1 mol Zn

1 mol Zn��65.4 g Zn

65.4 g Zn��1 mol Zn

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Section 10.3 Moles of CompoundsIn your textbook, read about chemical formulas and the mole, the molar mass of com-pounds, and conversions among mass, moles, and number of particles.

Study the table and the diagram of a methane molecule and a trichloromethane molecule. Then answer the following questions.

1. What elements and how many atoms of each does a molecule of methane contain?

2. What elements and how many atoms of each does a molecule of trichloromethane contain?

3. How many moles of each element are in a mole of methane?

4. How many moles of each element are in a mole of trichloromethane?

5. Which of the following values represents the number of carbon atoms in one mole ofmethane? 6.02 � 1023; 12.0 � 1023; 18.1 � 1023; 24.1 � 1023

6. Which of the following values represents the number of chlorine atoms in one mole oftrichloromethane? 6.02 � 1023; 1.20 � 1024; 1.81 � 1024; 2.41 � 1023

7. Which of the following values represents the molar mass of methane? 13.02 g/mol; 16.05 g/mol; 52.08 g/mol; 119.37 g/mol

8. Chloromethane (CH3Cl) has a molar mass of 50.49 g/mol. Which of the following valuesrepresents the number of molecules of CH3Cl in 101 grams of the substance? 3.01 � 1023; 6.02 � 1023; 1.20 � 1024; 6.08 � 1026


Element Molar Mass (g/mol)

Hydrogen 1.01

Carbon 12.01

Chlorine 35.45

HC

H

HCH4

C

H

Cl

Cl

Cl

CHCl3

H

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Section 10.4 Empirical and Molecular FormulasIn your textbook, read about percent composition.


1. What is the percent composition of a compound?

2. Describe how to find the percent composition of a compound if you know the mass of asample of a compound and the mass of each element in the sample.

In your textbook, read about empirical and molecular formulas.

Circle the letter of the choice that best answers the question.

3. Which information about a compound can you use to begin to determine the empiricaland molecular formulas of the compound?

a. mass of the compound c. percent composition of the compound

b. number of elements in the compound d. volume of the compound

4. You have determined that a compound is composed of 0.300 moles of carbon and 0.600 moles of oxygen. What must you do to determine the mole ratio of the elements inthe empirical formula of the compound?

a. Multiply each mole value by 0.300 mol. c. Divide each mole value by 0.300 mol.

b. Multiply each mole value by 0.600 mol. d. Divide each mole value by 0.600 mol.

5. The mole ratio of carbon to hydrogen to oxygen in a compound is 1 mol C : 2 mol H : 1 mol O. What is the empirical formula of the compound?

a. CHO b. CH2O c. C2HO2 d. C2H2O2

6. You calculate the mole ratio of oxygen to aluminum in a compound to be 1.5 mol O : 1 mol Al. What should you do to determine the mole ratio in the empirical formula of the compound?

a. Multiply each mole value by 1.5. c. Divide each mole value by 1.5.

b. Multiply each mole value by 2. d. Divide each mole value by 2.

7. What is the relationship between the molecular formula and the empirical formula of acompound?

a. (molecular formula)(empirical formula) � n

b. molecular formula �

c. molecular formula � (empirical formula)n

d. molecular formula �n

��empirical formula

empirical formula��n


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8. You know that the empirical formula of a compound has a molar mass of 30.0 g/mol.The experimental molar mass of this compound is 60.0 g/mol. What must you do todetermine the value of n in the relationship between the molecular formula and theempirical formula?

a. Add 30.0 g/mol and 60.0 g/mol. c. Divide 60.0 g/mol by 30.0 g/mol.

b. Divide 30.0 g/mol by 60.0 g/mol. d. Multiply 30.0 g/mol by 60.0 g/mol.

9. You know that the experimental molar mass of a compound is three times the molar massof its empirical formula. If the compound’s empirical formula is NO2, what is its molecu-lar formula?

a. NO2 b. NO6 c. N3O2 d. N3O6

Solve the following problem. Show your work in the space provided.

10. A sample of a compound contains 7.89 g potassium, 2.42 g carbon, and 9.69 g oxygen.Determine the empirical and molecular formulas of this compound, which has a molarmass of 198.22 g/mol.



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Section 10.5 The Formula for a HydrateIn your textbook, read about naming and analyzing hydrates.

Use each of the terms below just once to complete the passage.

A(n) (1) is a compound that has a specific number of water

molecules bound to its atoms. Molecules of water that become part of a hydrate are called

waters of (2) . In the formula for a hydrate, the number of

(3) associated with each (4) of the

compound is written following a dot.

The substance remaining after a hydrate has been heated and its waters of hydration

released is called (5) . The ratio of the number of moles of

(6) to one mole of the anhydrous compound indicates the

coefficient of H2O that follows the dot in the formula of the hydrate. Because the anhydrous

form of the hydrate can absorb water into its (7) , hydrates are used

as (8) , which are drying agents.

Complete the table of hydrates.

Solve the following problem. Show your work in the space provided.

11. A 2.00-g sample of a hydrate of iron(II) chloride produces 1.27 g of anhydrous iron(II) chlo-ride (FeCl2) after heating. Determine the empirical formula and the name of the hydrate.


anhydrous crystal structure desiccants formula unit

hydrate hydration water molecules water of hydration

Chemical Formula Name

CdSO4 Cadmium sulfate, anhydrous

CdSO4�H20 9.

10. Cadmium sulfate tetrahydrate

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CHAPTER 10

Assessment



1. 4. 7. 10.

2. 5. 8.

3. 6. 9.

Short Answer


11.

Extended Response


12.

SAT Subject Test:Chemistry

13. 15.

14. 16.

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Chapter 11 StoichiometryMiniLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 58

ChemLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 59


Math Skills Transparency Masters and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . . . 66

Study Guide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 72



Table ofContents

57

Reproducible Pages

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mini LAB 11Apply Stoichiometry

How much sodium carbonate (Na2CO3) is produced when baking soda decomposes?Baking soda is used in many baking recipes because it makes batter rise, which results ina light and fluffy texture. This occurs because baking soda, sodium hydrogen carbonate(NaHCO3), decomposes upon heating to form carbon dioxide gas according to the following equation.

2NaHCO3 0 Na2CO3 � CO2 � H2O

Materials ring stand, ring, clay triangle, crucible, crucible tongs, Bunsen burner, balance, 3.0 g baking soda (NaHCO3)

Procedure 1. Read and complete the lab safety form.

2. Create a data table to record your experimental data and observation.

3. Use a balance to measure the mass of a clean, dry crucible. Add about 3.0 g of sodiumhydrogen carbonate (NaHCO3), and measure the combined mass of the crucible andNaHCO3. Record both masses in your data table, and calculate the mass of the NaHCO3.

4. Use this starting mass of NaHCO3 and the balanced chemical equation to calculate themass of NaHCO3 that will be produced.

5. Set up a ring stand with a ring and clay triangle for heating the crucible.

6. Heat the crucible with a Bunsen burner, slowly at first and then with a stronger flame, for7–8 min. Record your observations during the heating.

7. Turn off the burner, and use crucible tongs to remove the hot crucible.WARNING: Do not touch the hot crucible with your hands.

8. Allow the crucible to cool, and then measure the mass of the crucible and NaHCO3.

Analysis

1. Describe what you observed during the heating of the baking soda.

2. Compare your calculated mass of NaHCO3 with the actual mass you obtained fromthe experiment.

3. Calculate Assume that the mass of Na2HCO3 that you calculated in Step 4 is theaccepted value for the mass of product that will form. Calculate the error and percenterror associated with the experimentally measured mass.

4. Identify sources of error in the procedure that led to errors calculated in Question 3.

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CHEMLAB 11

Safety Precautions• Always wear safety glasses and a lab apron.• Hot objects will not appear to be hot.• Do not heat broken, chipped, or cracked glassware.• Turn off the hot plate when not in use.

ProblemWhich reactant is the limit-ing reactant? How does theexperimental mole ratio ofFe to Cu compare with themole ratio in the balancedchemical equation? What isthe percent yield?

Objectives• Observe a single replace-

ment reaction. • Measure the masses of

iron and copper.• Calculate the moles of

each metal and the moleratio.

Materialsiron metal filings,

20 mesh copper(II) sulfate

pentahydrate(CuSO4�5H2O)

distilled waterstirring rod150-mL beaker

400-mL beaker100-mL graduated

cylinderweighing paperbalancehot platebeaker tongs

Determine the Mole RatioIron reacts with copper(II) sulfate in a single replacement reaction.

By measuring the mass of iron that reacts and the mass of coppermetal produced, you can calculate the ratio of moles of reactant tomoles of product. This mole ratio can be compared to the ratio foundin the balanced chemical equation.

Pre-Lab


2. Prepare all written materials that you will takeinto the laboratory. Be sure to include safety precautions and procedure notes. Use the datatable on the next page.

3. Is it important that you know you are using the hy-drated form of copper(II) sulfate? Would it be poss-ible to use the anhydrous form? Why or why not?

Procedure


2. Measure and record the mass of a clean, dry 150-mL beaker.

3. Place approximately 12 g of copper(II) sulfatepentahydrate into the 150-mL beaker and measureand record the combined mass.

4. Add 50 mL of distilled water to the copper(II)sulfate pentahydrate and heat the mixture on thehot plate at a medium setting. Stir until all of thesolid is dissolved, but do not boil. Using tongs,remove the beaker from the hot plate.

5. Measure approximately 2 g of iron metal filingsonto a piece of weighing paper. Measure andrecord the exact mass of the filings.

6. While stirring, slowly add the iron filings to thehot copper(II) sulfate solution.

7. Allow the reaction mixture to stand, without stir-ring, for 5 minutes to ensure complete reaction.The solid copper metal will settle to the bottomof the beaker.

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8. Use the stirring rod to decant (pour off) the liq-uid into a 400-mL beaker. Be careful to decantonly the liquid.

9. Add 15 mL of distilled water to the copper solidand carefully swirl the beaker to wash the cop-per. Decant the liquid into the 400-mL beaker.

10. Repeat step 9 two more times.

11. Place the 150-mL beaker containing the wet copper on the hot plate. Use low heat to dry thecopper.

12. Remove the beaker from the hot plate and allowit to cool.

13. Measure and record the mass of the cooled150-mL beaker and the copper.


1. Make sure the hot plate is off.

2. The dry copper can be placed in a waste con-tainer. Wet any residue that sticks to the beakerand wipe it out using a paper towel. Pour theunreacted copper(II) sulfate and iron(II) sulfatesolutions into a large beaker in the fume hood.

3. Return all lab equipment to its proper place.

4. Wash your hands thoroughly after all lab workand cleanup is complete.

CHEMLAB 11


1. Apply Write a balanced chemical equation for the reaction and calculate the mass ofcopper (Cu) that should have formed from the sample of iron (Fe) used. This mass is the theoretical yield.

2. Interpret Data Using your data, determine the mass and the moles of copper produced. Calculate the moles of iron used, and determine the whole-number iron-to-copper mole ratio and percent yield.

Mass of empty 150-mL beaker

Mass of 150-mL beaker � CuSO4�5H2O

Mass of CuSO4�5H2O

Mass of iron filings

Mass of 150-mL beaker and dried copper

Mass of dried copper

Observations

Data for the Reaction of Copper(II) Sulfate and Iron

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3. Compare and Contrast Compare the theoretical iron-to-copper mole ratio to the mole ratioyou calculated using the experimental data.

4. Error Analysis Identify sources of the error that resulted in deviation from the mole ratiogiven in the balanced chemical equation.

Inquiry Extension

Compare your results with those of several other lab teams. Create a hypothesis to explainany differences.

CHEMLAB 11

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Mole of given substance

Mass of given substance Mass of unknown substance

Moles of unknown substance

Step 1Start with a

balanced equation

Step 2Convert

from grams to moles

Step 4Convert

from moles to grams

Step 3Convert from molesof given to moles of

unknown

moles of unknownmoles of given

1 m

ol

nu

mb

er o

f g

ram

s

nu

mb

er o

f g

ram

s1

mo

l

no direct conversion

Mass-to-Mass ConversionsMass-to-Mass Conversions



35

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1. What conversion factor would you use to convert correctly from the mass of a given substance to the number of moles of the given substance?

2. What conversion factor would you use to convert correctly from the number of moles ofa given substance to the number of moles of an unknown substance?

3. What conversion factor would you use to convert correctly from the number of moles ofthe unknown substance to the mass of the unknown substance?

4. What is the name of the conversion factor in question 2?

5. What do you need to know to find the conversion factor in question 2?

Use the following balanced chemical equation and table to answer questions 6.

2N2(g) � O2(g) 0 2N2O(g)

6. Write the conversion factors in the order you would use them to determine correctly eachof the following.

a. the number of moles of N2O produced when 26.5 g N2 reacts with excess oxygen

b. the mass of N2 needed to produce 11.5 g N2O

c. the mass of N2 needed to react completely with 1.56 g O2

d. the mass of N2O produced when 7.05 g O2 reacts with excess nitrogen

Mass-to-Mass ConversionsMass-to-Mass Conversions



35

Compound Molar Mass (g/mol)

N2 28.02

O2 32.00

N2O 44.02

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3 N

itro

gen

mo

lecu

les

6 N

itro

gen

ato

ms

Bef

ore

rea

ctio

nA

fter

rea

ctio

n

3 H

ydro

gen

mo

lecu

les

6 H

ydro

gen

ato

ms

2 A

mm

on

ia m

ole

cule

s2

Nit

rog

en m

ole

cule

s4

Nit

rog

en a

tom

s

��

Limiting ReactantsLimiting Reactants



36

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1. How many N2 molecules are shown in the transparency? N atoms?

2. How many H2 molecules are shown? H atoms?

3. What is the ratio of H atoms to N atoms in one NH3 molecule?

4. How many H atoms would be needed to react with all the N atoms shown in the transparency?

5. How many N atoms would be needed to react with all the H atoms shown in the transparency?

6. According to your answers to questions 4 and 5, how many N2 molecules and H2molecules will be used up completely by the reaction shown in the transparency?

7. Which reactant will remain after the reaction? How many molecules?

8. Complete the diagram below by drawing the products of the chemical reaction in the box.

9. Which reactant in the diagram is the limiting reactant?

10. Which reactant in the diagram is in excess?

Reactant 1 Reactant 2

Before reaction

�

After reaction

Product

Limiting ReactantsLimiting Reactants



36

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Wh

at m

ole

rat

ios

can

be

der

ived

fro

m t

he

follo

win

g b

alan

ced

eq

uat

ion

?

4P(s

) �

5O

2(g

) 0

1P 4

O10

(s)

4 m

ol P

� 5

mo

l O2 0

1 m

ol P

4O10

4 m

ol P

4 m

ol P

� 5

mo

l O2

4 m

ol P

0 1

mo

l P4O

10

4 m

ol P

4 m

ol P

1 m

ol P

4O10

�

5 m

ol O

2

1 m

ol P

4O10

0 1

mo

l P4O

10

1 m

ol P

4O10

4 m

ol P

5 m

ol O

2�

5 m

ol O

2

5 m

ol O

2

0 1

mo

l P4O

10

5 m

ol O

2

5 m

ol O

2

4 m

ol P

1 m

ol P

4O10

4 m

ol P

1 m

ol P

4O10

5 m

ol O

2

4 m

ol P

5 m

ol O

2

4 m

ol P

1 m

ol P

4O10

5 m

ol O

2

1 m

ol P

4O10

1

1

1

Determining Mole RatiosDetermining Mole Ratios



15

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Determine the mole ratios for each of the following balanced chemical equations.

1. 2C(s) � O2(g) 0 2CO(g)

2. WO3(s) � 3H2(g) 0 W(s) � 3H2O(g)

3. 2IrCl3(aq) � 3NaOH(aq) 0 Ir2O3(s) � 3HCl(aq) � 3NaCl(aq)

Determining Mole RatiosDetermining Mole Ratios



15

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Mole ratios relate moles of unknown and known substances in a balanced chemical equation and are used to make stoichiometric mole-to-mole conversions between molar amounts of unknown and known substances.

To convert

number of moles of known substance

number of moles of unknown substance mole ratio

to usethis

1 mol C3H8

5 mol O2

1 mol C3H8

3 mol CO2

1 mol C3H8

4 mol H2O

5 mol O2

1 mol C3H8

5 mol O2

3 mol CO2

5 mol O2

4 mol H2O

3 mol CO2

1 mol C3H8

3 mol CO2

5 mol O2

3 mol CO2

4 mol H2O

4 mol H2O

1 mol C3H8

4 mol H2O

5 mol O2

4 mol H2O

3 mol CO2

C3H8 O2

C3H8 CO2

C3H8 H2O

O2 C3H8

O2 CO2

O2 H2O

CO2 C3H8

CO2 O2

CO2 H2O

H2O C3H8

H2O O2

H2O CO2

C3H8(g) � 5O2(g) 0 3CO2(g) � 4H2O(g)

Using Mole RatiosUsing Mole Ratios



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For each of the following problems, write the balanced chemical equation that represents the reaction. Then complete the table below by identifying the known substance, the unknown substance, and the mole ratio that you would use to solve each problem correctly.

1. Copper(II) oxide (CuO) decomposes into copper (Cu) and oxygen (O2) gas. What massof copper will be produced by the decomposition of 1.25 kg CuO?

2. Ammonia (NH3) is produced by the reaction of nitrogen (N2) and hydrogen (H2) gases.How much ammonia will be produced if 22.0 g H2 reacts with excess N2?

3. The reaction of sodium (Na) and water (H2O) produces sodium hydroxide (NaOH) andhydrogen (H2) gas. What mass of hydrogen gas is produced if 17.54 g NaOH is producedby the reaction?

4. The combustion of acetic acid (HC2H3O2) produces carbon dioxide (CO2) and water(H2O). What mass of carbon dioxide will be produced from the combustion of 25.0 gHC2H3O2?

5. 20.0 g of iron(III) sulfide (Fe2S3) was prepared by heating iron (Fe) and excess sulfur(S). What mass of iron was used in the preparation?

Using Mole RatiosUsing Mole Ratios



16

Chemical Formula Chemical FormulaProblem of Known Substance of Unknown Substance Mole Ratio

1

2

3

4

5

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What mass in grams of silver (Ag) will be produced when 125 g of silver oxide (Ag2O) decomposes?

Write the balanced equation.

Determine the number of moles of the unknown substance from the number of moles of the known substance using mole-to-mole conversions.

Determine the mass of the unknown substance using mole-to-mass conversion.

2Ag2O 0 4Ag � O2

0.539 g Ag2O �

4 mol Ag

2 mol Ag2O

� 1.08 mol Ag

1.08 mol Ag � 107.87 g Ag

1 mol Ag� 116 g Ag

Step 1

Step 2

Step 3

Step 4

unknown substance

known substance

Find the number of moles of known substance using mass-to-mole conversion.

1 mol Ag2O � � 215.74 g Ag107.87 g Ag

1 mol Ag

2 mol Ag

1 mol Ag2O�

1 mol Ag2O � � 16.00 g O 16.00 g O

1 mol O

1 mol O

1 mol Ag2O�

125 g Ag2O � � 0.539 mol Ag2O1 mol Ag2O

231.74 g Ag2O

Molar mass Ag2O � 231.74 g Ag2O

number of moles ofknown substance

number of moles ofunknown substance

Solving Stoichiometric Mass-to-Mass Conversion ProblemsSolving Stoichiometric Mass-to-Mass Conversion Problems



17

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1. The reaction of iron(III) oxide (Fe2O3) and hydrogen (H2) is represented by the follow-ing unbalanced chemical equation.

Fe2O3(s) � H2(g) 0 Fe(s) � H2O(l)

Determine the mass in grams of hydrogen gas needed to react completely with 33.5 g Fe2O3.

Step 1.

Step 2.

Step 3.

Step 4.

2. Determine the mass in grams of copper(II) sulfide (Cu2S) formed when 15.0 g copper(I)chloride (CuCl) reacts with excess hydrogen sulfide (H2S) according to the followingunbalanced chemical equation.

CuCl(aq) � H2S(g) 0 Cu2S(s) � HCl(aq)

Step 1.

Step 2.

Step 3.

Step 4.

Solving Stoichiometric Mass-to-Mass Conversion ProblemsSolving Stoichiometric Mass-to-Mass Conversion Problems



17

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StoichiometryStoichiometry

Section 11.1 What is stoichiometry?In your textbook, read about stoichiometry and the balanced equation.


1. The study of the quantitative relationships between the amounts ofreactants used and the amounts of products formed by a chemical reactionis called stoichiometry.

2. Stoichiometry is based on the law of conservation of mass.

3. In any chemical reaction, the mass of the products is less than the mass ofthe reactants.

4. The coefficients in a chemical equation represent not only the number ofindividual particles but also the number of moles of particles.

5. The mass of each reactant and product is related to its coefficient in thebalanced chemical equation for the reaction by its molar mass.

Complete the table below, using information represented in the chemical equation forthe combustion of methanol, an alcohol.

methanol � oxygen 0 carbon dioxide � water

2CH3OH(l) � 3O2 (g) 0 2CO2(g) � 4H2O(g)

10. What are the reactants?

11. What are the products?

12. What is the total mass of the reactants?

13. What is the total mass of the products?

14. How do the total masses of the reactants and products compare?


Number of Number of Substance Molar Mass (g/mol) Molecules Moles (mol) Mass (g)

6. Methanol 32.05

7. Oxygen gas 32.00

8. Carbon dioxide 44.01

9. Water 18.02

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In your textbook, read about mole ratios.

Answer the questions about the following chemical reaction.

sodium � iron(III) oxide 0 sodium oxide � iron

6Na(s) � Fe2O3(s) 0 3Na2O(s) � 2Fe(s)

15. What is a mole ratio?

16. How is a mole ratio written?

17. Predict the number of mole ratios for this reaction.

18. What are the mole ratios for this reaction?

19. What is the mole ratio relating sodium to iron?

20. What is the mole ratio relating iron to sodium?

21. Which mole ratio has the largest value?


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Section 11.2 Stoichiometric CalculationsIn your textbook, read about mole-to-mole conversion.

Read the following passage and then solve the problems. In the equation that followseach problem, write in the space provided the mole ratio that can be used to solve theproblem. Complete the equation by writing the correct value on the line provided.

The reaction of sodium peroxide and water produces sodium hydroxide and oxygengas. The following balanced chemical equation represents the reaction.

2Na2O2(s) � 2H2O(l) 0 4NaOH(s) � O2(g)

1. How many moles of sodium hydroxide are produced when 1.00 mol sodium peroxidereacts with water?

1.00 mol Na2O2 � � mol NaOH

2. How many moles of oxygen gas are produced when 0.500 mol Na2O2 reacts with water?

0.500 mol Na2O2 � � mol O2

3. How many moles of sodium peroxide are needed to produce 1.00 mol sodium hydroxide?

1.00 mol NaOH � � mol Na2O2

4. How many moles of water are required to produce 2.15 mol oxygen gas in this reaction?

2.15 mol O2 � � mol H2O

5. How many moles of water are needed for 0.100 mol of sodium peroxide to react com-pletely in this reaction?

0.100 mol Na2O2 � � mol H2O

6. How many moles of oxygen are produced if the reaction produces 0.600 mol sodiumhydroxide?

0.600 mol NaOH � � mol O2

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In your textbook, read about mole-to-mass and mass-to-mass conversions.

Solving a mass-to-mass problem requires the four steps listed below. The equations in theboxes show how the four steps are used to solve an example problem. After you have stud-ied the example, solve the problems below, using the four steps.

Example problem: How many grams of carbon dioxide are produced when 20.0 g acetylene(C2H2) is burned?

Step 1 Write a balanced chemical equation for the reaction.

Step 2 Determine the number of moles of the known substance,using mass-to-mole conversion.

Step 3 Determine the number of moles of the unknown substance,using mole-to-mole conversion.

Step 4 Determine the mass of the unknown substance, using mole-to-mass conversion.

7. In some mole-to-mass conversions, the number of moles of the known substance is given. In those conversions, which step of the above solution is not necessary?

8. In a blast furnace, iron and carbon monoxide are produced from the reaction of iron(III)oxide (Fe2O3) and carbon. How many grams of iron are formed when 150 g iron(III) oxidereacts with an excess of carbon?

9. Solid sulfur tetrafluoride (SF4) and water react to form sulfur dioxide and an aqueous solu-tion of hydrogen fluoride. How many grams of water are necessary for 20.0 g sulfur tetra-fluoride to react completely?


Solution

2C2H2(g) � 5O2(g)

0 4CO2(g) � 2H2O(g)

20.0 g C2H2 �

� 0.768 mol C2H2

0.768 mol C2H2 �

� 1.54 mol CO2

1.54 mol CO2 �

� 67.8 g CO2

44.01 g CO2��1 mol CO2

4 mol CO2��2 mol C2H2

1 mol C2H2��26.04 g C2H2

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Section 11.3 Limiting ReactantsIn your textbook, read about why reactions stop and how to determine the limitingreactant.

Study the diagram showing a chemical reaction and the chemical equation that repre-sents the reaction. Then complete the table. Show your calculations for questions 25–27in the space below the table.

O2 � 2NO 0 2NO2

The molar masses of O2, NO, and NO2 are 32.00 g/mol, 30.01 g/mol, and 46.01 g/mol,respectively.

� � 0 �

Amount and Name Amount of O2 Amount of NO Amount of NO2 Limiting Reactant of Excess Reactant

1 molecule 2 molecules 2 molecules none none

4 molecules 4 molecules 4 molecules NO 2 molecules O2

2 molecules 8 molecules 1. 2. 3.

1.00 mol 2.00 mol 4. 5. 6.

4.00 mol 4.00 mol 7. 8. 9.

5.00 mol 7.00 mol 10. 11. 12.

1.00 mol 4.00 mol 13. 14. 15.

0.500 mol 0.200 mol 16. 17. 18.

32.00 g 60.02 g 19. 20. 21.

16.00 g 80.00 g 22. 23. 24.

10.00 g 20.00 g 25. 26. 27.

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Section 11.4 Percent YieldIn your textbook, read about the yields of products.

Study the diagram and the example problem.

Example Problem: The following chemical equation represents the production of gallium oxide,a substance used in the manufacturing of some semiconductor devices.

4Ga(s) � 3O2(g) 0 2Ga2O3(s)

In one experiment, the reaction yielded 7.42 g of the oxide from a 7.00-g sample of gallium.Determine the percent yield of this reaction. The molar masses of Ga and Ga2O3 are 69.72 g/mol and 187.44 g/mol, respectively.

Use the information in the diagram and example problem to evaluate each value orexpression below. If the value or expression is correct, write correct. If it is incorrect, writethe correct value or expression.

1. actual yield: unknown

2. mass of reactant: 7.00 g Ga

3. number of moles of reactant: 7.00 g Ga �

4. number of moles of product: 0.100 mol Ga �

5. theoretical yield: 0.0500 mol Ga2O3 �

6. percent yield: � 100 9.37 g Ga2O3��7.42 g Ga2O3

187.44 g Ga2O3��1 mol Ga2O3

2 mol Ga2O3��1 mol Ga

69.72 g Ga��1 mol Ga


percent yield � � 100%actual yield

��theoretical yield

mass of product from experimentalmeasurement

mass of product predicted from stoichiometric calculation using

a. mass of reactantb. 4-step mass-to-mass conversion

1. Write the balanced chemical equation.2. Calculate the number of moles of reactant, using

molar mass.3. Calculate the number of moles of product, using

the appropriate mole ratio.4. Calculate the mass of product, using the reciprocal

of molar mass.

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1. 4. 7.

2. 5. 8.

3. 6.

Short Answer


9.

Extended Response


10.

11.

SAT Subject Test: Chemistry

12. 14. 16.

13. 15. 17.

CHAPTER 11

Assessment

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Chapter 12 States of MatterMiniLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 86

ChemLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 87

Teaching Transparency Masters and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . . 90

Math Skills Transparency Masters and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . . 94

Study Guide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 96

Chapter Assessment. . . . . . . . . . . . . . . . . . . . . . . . . . . . 102

STP Recording Sheet . . . . . . . . . . . . . . . . . . . . . . . . . . . 108

Table ofContents

85

Reproducible Pages

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mini LAB 12Model Crystal Unit Cells

Formulating Models You can make physical models that illustrate the structures of crystals.

Materials plastic or paper soda straws (12), 22- or 26-gauge wire, scissors

Procedure 1. Read and complete the lab safety form.

2. Cut four soda straws into thirds. Wire the pieces to make a cube. All angles are 90°.

3. To model a rhombohedral crystal, deform the cube from step 1 until no angles are 90°.

4. To model a hexagonal crystal, flatten the model from step 2 until it looks like a piewith six slices.

5. To model a tetragonal crystal, cut four straws in half. Cut four of the pieces in halfagain. Wire the eight shorter pieces to make four square ends. Use the longer piecesto connect the square ends.

6. To model the orthorhombic crystal, cut four straws in half. Cut 1/3 off four of thehalves. Connect the four long, four medium, and four short pieces so that each side is a rectangle.

7. To model the monoclinic crystal, deform the model from step 5 along one axis. Tomodel the triclinic crystal, deform the model from step 5 until it has no 90° angles.

Analysis

1. Evaluate Which two models have three axes of equal length? How do these modelsdiffer?

2. Determine which model includes a square and rectangle.

3. Determine which models have three unequal axes.

4. Infer Do you think crystals are perfect or do they have defects? Explain your answer.

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CHEMLAB 12

Safety Precautions• Always wear safety goggles and a lab apron.• Wear gloves because some of the liquids can dry out your skin. • Avoid inhaling any of the vapors, especially ammonia. • There should be no open flames in the lab; some of the liquids are

flammable.

ProblemHow do intermolecularforces affect the evapora-tion rates of liquids?

Objectives• Measure and compare the

rates of evaporation fordifferent liquids.

• Classify liquids based ontheir rates of evaporation.

• Predict which intermolec-ular forces exist betweenthe particles of each liquid.

Materialsdistilled waterethanolisopropyl alcoholacetonehousehold

ammoniadroppers (5)small plastic cups (5)

grease pencil ormasking tape anda marking pen

paper towelsquare of waxed

paperstopwatch

Compare Rates of EvaporationSeveral factors determine how fast a sample of liquid will evaporate.

The volume of the sample is a key factor. A drop of water takes lesstime to evaporate than a liter of water. The amount of energy suppliedto the sample is another factor. In this lab, you will investigate how thetype of liquid and temperature affect the rate of evaporation.

Pre-Lab

1. Read the entire CHEMLAB. Use the data tableon the next page.

2. What is evaporation? Describe what happens atthe molecular level during evaporation.

3. List the three possible intermolecular forces. Whichforce is the weakest? Which force is the strongest?

4. Look at the materials list for this lab. Consider thefive liquids you will test. Predict which liquidswill evaporate quickly and which will take longerto evaporate. Give reasons for your predictions.

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5. To calculate an evaporation rate, you would dividethe evaporation time by the quantity of liquidused. Explain why it is possible to use the evapo-ration times from this lab as evaporation rates.

6. Make sure you know how to use the stopwatchprovided. Will you need to convert the reading onthe stopwatch to seconds?

Procedure


2. Use a grease pencil or masking tape to label eachof five small plastic cups. Use A for distilledwater, B for ethanol, C for isopropyl alcohol, Dfor acetone, and E for household ammonia.

3. Place the plastic cups on a paper towel.

4. Use a dropper to collect about 1 mL of distilledwater and place the water in the cup labeled A.Place the dropper on the paper towel directly infront of the cup. Repeat with the other liquids.

5. Place a square of waxed paper on your lab sur-face. Plan where on the waxed paper you will

place each of the 5 drops that you will test. Thedrops must be as far apart as possible to avoid mixing.

6. Have your stopwatch ready. Collect some waterin your water dropper and place a single drop onthe waxed paper. Begin timing. Time how long ittakes for the drop to completely evaporate. Whileyou wait, make two drawings of the drop. Onedrawing should show the shape of the drop asviewed from above. The other drawing should bea side view at eye level. If the drop takes longerthan 5 minutes to evaporate, record � 300 in yourdata table.

7. Repeat step 6 with the four other liquids.

8. Use the above procedure to design an experimentin which you can observe the effect of tempera-ture on the rate of evaporation of ethanol. Yourteacher will provide a sample of warm ethanol.Record your observations.


1. Crumple up the waxed paper and place it in thecontainer assigned by your teacher.

2. Place unused liquids in the containers specifiedby your teacher.

3. Wash out all droppers and test tubes except thoseused for distilled water.

CHEMLAB 12


1. Classifying Which liquids evaporated quickly? Which liquids were slow to evaporate?

Liquid Evaporation time (s) Shape of liquid drop

Distilled water

Ethanol

Ethanol (warm)

Isopropyl alcohol

Acetone

Household ammonia

Evaporation Data

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2. Conclude Based on your data, in which liquid(s) are the attractive forces betweenmolecules most likely to be dispersion forces?

3. Interpret Data Make a generalization about the shape of a liquid drop and theevaporation rate of the liquid.

4. Recognize Cause and Effect What is the relationship between surface tension and theshape of a liquid drop? What are the attractive forces that increase surface tension?

5. Apply Concepts The isopropyl alcohol you used is a mixture of isopropyl alcohol andwater. Would pure isopropyl alcohol evaporate more quickly or more slowly compared tothe alcohol and water mixture? Give a reason for your answer.

6. Think Critically Household ammonia is a mixture of ammonia and water. Based on thedata you collected, is there more ammonia or more water in the mixture? Use what youlearned about the relative strengths of the attractive forces in ammonia and water tosupport your conclusion.

7. Conclude How does the rate of evaporation of warm ethanol compare to ethanol at roomtemperature? Use kinetic-molecular theory to explain your observations.

8. How could you change the procedure to make it more precise?

Inquiry Extension

1. The vapor phases of liquids such as acetone and alcohol are more flammable than theirliquid phases. For flammable liquids, what is the relationship between evaporation rate andthe likelihood that the liquid will burn?

2. Suggest why a person who has a higher than normal temperature might be given arubdown with rubbing alcohol (70% isopropyl alcohol).

3. Table salt can be collected from salt water by evaporation. The water is placed in large,shallow containers. What advantage do these shallow containers have over deep containerswith the same overall volume?

Error Analysis

CHEMLAB 12

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Vac

uu

m

Vac

uu

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Clo

sed

end

Mer

cury

leve

lseq

ual

�h

ManometerManometer



37

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Teaching Transparency Worksheets Chemistry: Matter and Change • Chapter 12 91Teaching Transparency Worksheets Chemistry: Matter and Change • Chapter 12 91

1. The transparency shows a manometer. Briefly describe its parts.

2. Why are the levels of the mercury in the two arms of the U-tube the same when there isno gas in the flask?

3. What happens when gas enters the flask?

4. Compare the force exerted by the gas to the force exerted by the mercury contained inthe portion of the closed-end arm labeled �h.

5. If the flask was filled with air at a pressure of 1.00 atm, what would be the value of �hin millimeters of Hg?

6. If the flask was filled with air at a pressure of 0.50 atm, what would be the value of �hin millimeters of Hg?

7. A sample of gas is collected in the flask. The value of �h is 76.0 mm Hg. What is thepressure of the gas in mm Hg? In atm?

8. How does the vapor pressure of mercury affect the pressure reading of the manometer?

ManometerManometer



37

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1.0

�78 31

LIQUIDSOLID

GAS

Pres

sure

(at

m)

Phase Diagram for CO2

Temperature (C�)

100.00 373.990.00

Pres

sure

(at

m)

1.00

217.75Criticalpressure

Criticaltemperature

Temperature (C )

Phase Diagram for H2O

Normalfreezingpoint

Normalboilingpoint

LIQUID

Critical point

Triple point

A

B

VAPOR

SOLID

Triple point(�57°C,5.1 atm)

Criticalpoint(31°C,73 atm)

Phase DiagramsPhase Diagrams



38

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1. What variables are plotted on a phase diagram?

2. How many phases of water are represented in its phase diagram? What are they?

3. Use the phase diagram for water to complete the following table.

4. What phases of water coexist at each point along the red curve?

5. What two phase changes occur at each point along the yellow curve in the phase diagramfor water?

6. Look at the phase diagram for carbon dioxide. Above which pressure and temperature iscarbon dioxide unable to exist as a liquid?

7. At which pressure and temperature do the solid, liquid, and gaseous phases of carbondioxide coexist?

Phase DiagramsPhase Diagrams



38

Temperature (°C) Pressure (atm) Phase

200 1

�2 1

150 100

�2 0.001

30 0.8

1 liquid

100.00 vapor

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Unit Cells of CrystalsUnit Cells of Crystals



18

c

ba

c

ba

c

ba

c

ba

c

ba

c ba

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1. Explain what a crystalline solid is.

2. How many surfaces, or faces, do most crystal unit cells have? Which type of unit cell has adifferent number of faces? What is that number?

3. How many corners do most unit cells have? Which type of unit cell has a different number of corners? What is that number?

4. What do the letters a, b, and c in the transparency represent?

5. What do the symbols �, �, and � represent?

6. How many dimensions (length, width, depth) are needed to classify a unit cell?

7. How many faces are needed to determine the dimensions of a unit cell?

8. How many angle measurements are needed to classify a unit cell?

9. Identify the types of unit cells that have three equal dimensions.

10. Identify the types of unit cells that have equal angles.

11. How does the cubic unit cell differ from the rhombohedral unit cell?

12. Which unit cells meet the requirements a b and � �?

13. How does the triclinic unit cell differ from all the other unit cells?

Unit Cells of CrystalsUnit Cells of Crystals



18

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States of MatterStates of Matter

Section 12.1 GasesIn your textbook, read about the kinetic-molecular theory.

Complete each statement.

1. The kinetic molecular theory describes the behavior of gases in terms of particles in

.

2. The kinetic-molecular theory makes the following assumptions.

a. In a sample of a gas, the volume of the gas particles themselves is very

compared to the volume of the sample.

b. Because gas particles are far apart, there are no significant attractive or repulsive

between gas particles.

c. Gas particles are in constant and motion.

d. The collisions between gas particles are ; that is, no

energy is lost.

3. The kinetic energy of a particle is represented by the equation .

4. is a measure of the average kinetic energy of the particles in asample of matter.

In your textbook, read about explaining the behavior of gases.


5. Gases are less dense than solids because there is a lot of space between the particles of a gas.

6. The random motion of gas particles causes a gas to expand until it fills its container.

7. The density of a gas decreases as it is compressed.

8. A gas can flow into a space occupied by another gas.

9. The diffusion of a gas is caused by the random motion of the particles of the gas.

10. Lighter gas particles diffuse less rapidly than do heavier gas particles.

11. During effusion. a gas escapes through a tiny opening into a vacuum.

12. Graham’s law of effusion states that the rate of effusion for a gas isdirectly related to the square root of its molar mass.

STUDY GUIDE CHAPTER 12

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In your textbook, read about gas pressure.

Circle the letter of the choice that best completes the statement or answers the question.

13. Pressure is defined as force per unit

a. area. b. mass. c. time. d. volume.

14. What is an instrument designed to measure atmospheric pressure?

a. barometer b. manometer c. sphygmomanometer d. thermometer

15. The height of the liquid in a barometer is affected by all of the following EXCEPT the

a. altitude. c. density of the liquid in the column.

b. atmospheric pressure. d. diameter of the column tube.

16. The pressure of the gas in a manometer is directly related to which of the followingquantities?

a. height of the mercury column in the closed-end arm

b. height of the mercury column in the open-end arm

c. a b

d. a � b

17. One atmosphere is equal to a pressure of

a. 76 mm Hg. b. 101.3 kPa. c. 147 psi. d. 706 torr.

18. The partial pressure of a gas depends on all of the following EXCEPT the

a. concentration of the gas. c. size of the container.

b. identity of the gas. d. temperature of the gas.

19. The pressure of a sample of air in a manometer is 102.3 kPa. What is the partial pressure of nitrogen (N2) in the sample if the combined partial pressures of the other gases is 22.4 kPa?

a. 62.4 kPa b. 79.9 kPa c. 102.3 kPa. d. 124.7 kPa

Use the figure to answer the following questions.

20. What instrument is illustrated in the figure?

21. Who invented this instrument?

22. What are the two opposing forces that control the height of the mercury in the column?

23. What does it mean when the level of mercury rises in the column?


STUDY GUIDE CHAPTER 12

Vacuum

Atmosphericpressure

Pressure exertedby mercury

column

760 mm

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Section 12.2 Forces of AttractionIn your textbook, read about forces of attraction.


1. Ionic, metallic, and covalent bonds are examples of what type of forces?

2. Dispersion forces, dipole–dipole forces, and hydrogen bonds are examples of what type

of forces?

3. Describe dispersion forces.

4. Dispersion forces are greatest between what type of molecules?

5. Describe a permanent dipole.

6. Describe dipole–dipole forces.

7. Describe a hydrogen bond.

8. Identify each of the diagrams below as illustrating dipole–dipole forces, dispersion forces,or hydrogen bonds.

a. b. c.

9. Rank dipole–dipole forces, dispersion forces, and hydrogen bonds in order of increasing strength.

��

��

��

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Section 12.3 Liquids and SolidsIn your textbook, read about liquids and solids.

In the space at the left, write true if the statement is true; if the statement is false,change the italicized word or phrase to make it true.

1. The constant motion of the particles in a liquid causes the liquid totake the shape of its container.

2. At room temperature and one atmosphere of air pressure, thedensity of a liquid is much greater than that of its vapor.

3. Liquids are not easily compressed because their particles areloosely packed.

4. A liquid is less fluid than a gas because intramolecular attractionsinterfere with the ability of particles to flow past one another.

5. Liquids that have stronger intermolecular forces have higherviscosities than do liquids with weaker intermolecular forces.

6. The viscosity of a liquid increases with temperature because theincreased average kinetic energy of the particles makes it easierfor the particles to flow.

7. Liquids that can form hydrogen bonds generally have a highsurface tension.

8. A liquid that rises in a narrow glass tube shows that the adhesiveforces between the particles of the liquid and glass are greaterthan the cohesive forces between the particles of the liquid.

9. Solids have a definite shape and volume because the motion oftheir particles is limited to vibrations around fixed locations.

10. Most solids are less dense than liquids because the particles in asolid are more closely packed than those in a liquid.

11. Rubber is a crystalline solid because its particles are not arrangedin a regular, repeating pattern.


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Section 12.4 Phase ChangesIn your textbook, read about phase changes.

Complete the table by writing the initial and final phases for each phase change andmaking a check (✔) in the correct energy column.

For each item in Column A, write the letter of the matching item in Column B.

Column A Column B

7. Temperature at which a liquid is converted into acrystalline solid

8. Temperature at which the forces holding a crystallinelattice together are broken

9. Temperature at which the vapor pressure of a liquidequals the external or atmospheric pressure


a. boiling point

b. freezing point

c. melting point

Phase ChangePhase Energy

initial final required released

1. Condensation

2. Deposition

3. Freezing

4. Melting

5. Sublimation

6. Vaporization

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In your textbook, read about phase diagrams.

Use the phase diagram for water to answer the following questions.

10. What variables are plotted on a phase diagram?

11. What phase of water is represented by each of the following regions?

a. Region I

b. Region II

c. Region III

12. What does point 2 represent?

13. What is the temperature at point 3?

14. What does line A represent?

15. What is point 4 called? What does it represent?

100.00 373.990.00

Pres

sure

(at

m)

1.00

217.75

Temperature (C�)

Region IIIB

4

3

1

2

C

A

Region II

Region I



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Name Date Class

CHAPTER 12

Assessment



1. 3. 5. 7.

2. 4. 6. 8.

Short Answer


9.

10.

Extended Response


11.

SAT Subject Test:Chemistry

12. 14.

13. 15.


chapters 9–12 resources - destiny high school · chapters 9-12 resources ... 4 chemistry: matter...

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