chapters 13–16 resources - wikispacesdearbornchemistry.wikispaces.com/file/view/cmcff13-16.pdf ·...

Chapters 13–16 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.

Send all inquiries to:Glencoe/McGraw-Hill8787 Orion PlaceColumbus, OH 43240-4027

ISBN: 978-0-07-878763-8MHID: 0-07-878763-7

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 13-16 Resources

Reproducible Student Pages

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

Chapter 13

Gases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1

Chapter 14

Mixtures and Solutions . . . . . . . . . . . . . . . . . . . . . . . . . 31

Chapter 15

Energy and Chemical Change . . . . . . . . . . . . . . . . . . . . . . 57

Chapter 16

Reaction Rates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 87

Teacher Guide and Answers

Chapter 13 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 113

Chapter 14 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 117

Chapter 15 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 122

Chapter 16 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 127

Table ofContents

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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 13 GasesMinilab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2

Chem Lab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .3

Teaching TransparencyMaster and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . . . . .6

Maths Skill TransparencyMaster and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . . . . 12

Study Guide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18

Chapter Assessment. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 24

STP Recording Sheet . . . . . . . . . . . . . . . . . . . . . . . . . . . . 30

Table ofContents

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2 Chemistry: Matter and Change • Chapter 13 ChemLab and MiniLab Worksheets

mini LAB 13Model a Fire Extinguisher

Why is carbon dioxide used in fire extinguishers?

Materials masking tape, aluminum foil, metric ruler, beaker, candle, matches, ther-mometer, barometer or weather radio, baking soda (NaHCO3), vinegar (5% CH3COOH)

Procedure1. Read and complete the lab safety form.

2. Measure the temperature with a thermometer. Obtain the air pressure with a barometer or weather ratio. Record your observations.

3. Roll a 23-cm � 30-cm piece of aluminum foil into a cylinder that is 30 cm long androughly 6 cm in diameter. Tape the edges with masking tape.

4. Use matches to light a candle. WARNING: Run water over the extinguished matchbefore throwing it away. Keep hair and loose clothing away from the flame.

5. Place 30 g of baking soda (NaHCO3) in a large beaker. Add 40 mL of vinegar (5% CH3 COOH).

6. Quickly position the foil cylinder at about 45° up and away from the top of the can-dle flame. WARNING: Do not touch the end of the aluminum tube that is nearthe burning candle.

7. While the reaction in the beaker is actively producing carbon dioxide gas, carefullypour the gas, but not the liquid, out of the beaker and into the top of the foil tube.Record your observations.

Analysis

1. Apply Calculate the molar volume of carbon dioxide gas (CO2) at room temperatureand atmospheric pressure.

2. Calculate the room-temperature densities in grams per liter of carbon dioxide, oxygen,and nitrogen gases. Recall that you will need to calculate the molar mass of each gas inorder to calculate densities.

3. Interpret Do your observations and calculations support the use of carbon dioxide gas toextinguish fires? Explain.

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ChemLab and MiniLab Worksheets Chemistry: Matter and Change • Chapter 13 3

Safety Precautions • Read and observe all cautions listed on the aerosol can of

office equipment duster. • Do not have any open flames in the room.

CHEMLAB 13

ProblemHow much pressure isrequired to burst a kernelof popcorn?

Objectives• Collect and analyze data

hard or soft water• Calculate the pressure

needed to burst a kernelof popcorn

• Design steps to controlexperimental error

Materialspopcorn kernels

(18-20) vegetable oil (1.5

mL) wire gauzesquares (2)

Bunsen burner ring stand

small iron ring 10-mL graduated

cylinder250-mL beakerbeaker tongs balance distilled waterpaper towels

Determine Pressure inPopcorn KernelsWhen the water vapor pressure inside a popcorn kernel

is great enough, the kernel bursts and releases thewater vapor. The ideal gas law can be used to find the pres-sure in the kernel as it bursts.

Pre-Lab

1. Read the entire CHEMLAB.

2. Water forms a meniscus when poured into a grad-uated cylinder. To correctly record the volume,which part of the meniscus do you read? Howmany digits do you write down when you recordthe measured volume?

3. In step 5, why is it important to dry the kernelsbefore moving on with the procedure?

4. How do you convert from mass of a compound tomoles of a compound? If you have 5.00 g ofwater, how many moles of water do you have?

5. Define atmospheric pressure (1atm) in terms ofkilopascals, millimeters of mercury, torrs, andbars.

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Procedure

1. Read and complete the lab safety form.

2. Create a table to record your data.

3. Place approximately 5 mL of distilled water inthe graduated cylinder, and record the volume.

4. Place 18-20 popcorn cornels in the graduatedcylinder with the water. Tap the cylinder toforce any air bubbles off the kernels. Record thenew volume.

5. Remove the kernels from the graduated cylinder, and dry them.

6. Place the dry kernels and 1.0-1.5mL of vegetable oil into the beaker.

7. Measure the total mass of the beaker, oil, andkernels.

8. Set up a Bunsen burner with a ring stand, ring,and wire guaze.

9. Place the beaker on the wire guaze and ring.Place another piece of wire guaze on top of thebeaker.

10. Gently heat the beaker with the burner. Movethe burner back and forth to heat the kernelsevenly.

11. Observe the changes in the kernels and oilwhile heating, then turn off the burner when thepopcorn has popped and before any burningoccurs.

12. Using the beaker tongs, remove the beaker fromthe ring and allow it to cool completely.

13. Measure the final mass of the beaker, oil, andpopcorn once cooling is complete.

14. Post your data at glencoe.com

15. Cleanup and Disposal Dispose of the popcornand oil as directed by your teacher. Wash andreturn all lab equipment to its designated location.

Analyze and Conclude

1. Calculate the volume of the popcorn kernels, in liters, by the difference in the volumes of distilledwater before and after adding popcorn.

2. Calculate the total mass of water vapor released using the mass measurements of the beaker, oil, andpopcorn before and after popping.

3. Convert Use the molar mass of water and the volume of popcorn to find the number of moles ofwater released.

4. Use Formulas Use the temperature of the boiling oil (225°C) as your gas temperature, and calculatethe pressure of the gas using the ideal gas law.

CHEMLAB 13

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5. Compare and Contrast atmospheric pressure to the pressure of the water vapor in thekernels.

6. Infer why all the popcorn kernels did not pop.

7. Error Analysis Identify a potential source of error for this lab, and suggest a method tocorrect it.

Inquiry Extension

Design an experiment that tests the amount of pressure necessary to different types ofpopcorn kernels.

CHEMLAB 13

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6 Chemistry: Matter and Change • Chapter 13 Teaching Transparency Masters

Volume (L)

Pressure vs. Volume for a Gas at Constant Temperature

Pres

sure

(kP

a)

250

200

150

100

50

010 2 3 4 5

(V1, P1)

(V2, P2)

Pressure vs. Volume GraphPressure vs. Volume Graph

TEACHING TRANSPARENCY MASTER

Use with Chapter 13,Section 13.1

39

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Teaching Transparency Worksheets Chemistry: Matter and Change • Chapter 13 7

1. Based on this graph, how is the volume of a gas affected by increased pressure at constant temperature?

2. The relationship between the volume and pressure of a gas at constant temperature is an inversely proportional relationship. Based on the evidence in this graph, define theinversely proportional relationship.

3. What do you notice when you multiply the pressure by the volume for any point on theline on the graph?

4. Based on the mathematical relationship derived from this graph, what is the pressure ofthe gas at 3.00 L at constant temperature?

5. What gas law does this graph represent?

6. What mathematical expression is used to define this law? Define all symbols used.

7. A sample of gas is compressed from 3.25 L to 1.20 L at constant temperature. If the pres-sure of this gas in the 3.25-L volume is 100.00 kPa, what will the pressure be at 1.20 L?List all known and unknown variables. Show all your work.

Pressure vs. Volume GraphPressure vs. Volume Graph

TEACHING TRANSPARENCY WORKSHEET


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Temperature (K)

Volume vs. Kelvin Temperaturefor a Gas at Constant Pressure

Vo

lum

e (L

)

5

4

3

2

1

02000 400 600 800 1000

(T1, V1)

(T2, V2)

Volume vs. Temperature GraphVolume vs. Temperature Graph



40

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1. Notice that this graph shows kelvin temperatures. How are the kelvin scale and Celsiusscale related mathematically?

2. Based on this graph, how is the volume of a gas affected by increased temperature atconstant pressure?

3. The relationship between the volume and temperature of a gas at constant pressure is adirectly proportional relationship. Based on the evidence in this graph, define the directlyproportional relationship.

4. What do you notice when you divide the temperature by the volume for any point on theline on the graph?

5. What law does this graph represent?

6. What mathematical expression is used to define this law? Define all symbols used.

7. The kelvin temperature of a sample of gas is decreased from 460 K to 240 K at constantpressure. If the volume of this gas at 460 K is 2.50 L, what will the volume be at 240 K?List all known and unknown variables. Show all your work.

Volume vs. Temperature GraphVolume vs. Temperature Graph



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2 m

ol

2 vo

lum

es2

mo

l2

volu

mes

1 m

ol

1 vo

lum

e1

mo

l1

volu

me

� ��

0 0M

eth

ane

gas

CH

4 (g)

Oxy

gen

gas

2O2 (g

)C

arb

on

dio

xid

e g

asC

O2 (g

)W

ater

vap

or

2H

2O(g

)

TEACHING TRANSPARENCY MASTER 41

Burning of Methane GasBurning of Methane Gas Use with Chapter 13,Section 13.3

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1. What do the coefficients in chemical equations involving gases, solids, and liquids represent?

2. What does Avogadro’s principle state?

3. Based on Avogadro’s principle, the coefficients in chemical equations involving onlygases represent two types of quantities. Name the two quantities.

4. Based on the balanced equation for the complete combustion of methane, how manyliters of carbon dioxide, CO2(g), and water vapor, H2O(g), are produced by the completecombustion of 1 L of methane gas, CH4?

5. What volume of oxygen gas is needed for the complete combustion of 8.00 L of methanegas, CH4? Assume that the pressure and temperature of the reactants are the same. Showall your work.

6. Write a balanced equation for the complete combustion of propane gas, C3H8, with oxygen, O2, to form carbon dioxide, CO2, and water, H2O.

7. What volume of carbon dioxide gas, CO2, is produced when 7.00 L of propane gas,C3H8, undergoes complete combustion, as shown in your answer to question 6? Show allyour work.

Burning of Methane GasBurning of Methane Gas



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12 Chemistry: Matter and Change • Chapter 13 Math Skills Transparency Masters

Solving Gas Problems Using Boyle’s Law

MATH SKILLS TRANSPARENCY MASTER


19

Solving Gas Problems Using Boyle’s Law

How can you solve gas problems using Boyle’s law?ProblemA sample of helium gas in a balloon has a pressure of 240 kPa in a 1.4-L container. What will be the new pressure if the gas sample is transferred to a 3.0-L container?

Step 1. Analyze the problem.

Known VariablesV1 � 1.4 L V2 � 3.0 L P1 � 240 kPa

Unknown VariableP2 � ? kPa

Step 2. Solve for the unknown.Divide both sides of the equation for Boyle’s law by V2 to solve for P2.

P1V1 � P2V2 P2 � P1� �Substitute the known values into the rearranged equation.

P2 � 240 kPa � �Multiply and divide numbers and units to solve for P2.

P2 � 240 kPa � � � 110 kPa

Step 3. Evaluate the answer.

1.4 L�3.0 L

1.4 L�3.0 L

V1�V2

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Math Skills Transparency Worsheets Chemistry: Matter and Change • Chapter 12 13

1. What two variables are inversely proportional, as stated in Boyle’s law?

2. What variable must be kept constant if you are solving a problem using Boyle’s law?

3. What is the first step in solving pressure-volume gas problems?

4. What is the equation for Boyle’s law?

5. How is the equation for Boyle’s law rearranged before solving it?

6. What is done to similar units in the numerator and the denominator in the final solving step?

7. What unit remains at the end? Is this the desired unit? Of what quantity is it a unit?

8. If the volume is almost doubled in a container at a fixed temperature, what will happento the pressure?

9. In step 3, how would you evaluate the answer to see whether or not it is reasonable?

10. If the known variables in a new problem are V1, P1, and P2, rewrite the equation forBoyle’s law to solve for the unknown variable.

Solving Gas Problems Using Boyle’s LawSolving Gas Problems Using Boyle’s Law

MATH SKILLS TRANSPARENCY WORKSHEET


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Solving Gas Problems Using the Combined Gas LawSolving Gas Problems Using the Combined Gas Law



20

How can you solve gas problems using the combined gas law?ProblemA gas at 100.0 kPa and 25.0°C fills a flexible container withan initial volume of 2.00 L. If the temperature is raised to60.0° and the pressure increased to 320.0 kPa, what is thenew volume?


Known VariablesP1 � 100.0 k Pa P2 � 320.0 kPaT1 � 25.0°C T2 � 60.0°C V1 � 2.00 LUnknown VariableV2 � ? L

Step 2. Solve for the unknown.Add 273 to the Celsius temperature for T1 and T2 to obtain the kelvin temperature.

T1 � 273 � 25.0°C � 298 K; T2 � 273 � 60.0°C � 333 KMultiply both sides of the equation for the combined law by T2 and divide by P2 to solve for V2.

� V2 � V1 � �� Substitute the known values into the rearranged equation; multiply and divide numbers and units to solve for V2.

V2 � 2.00 L� �� 0.698 L


333 K�298 K

100.0 kPa��320.0 kPa

T2�T1

P1�P2

P2V2�T2

P1V1�T1

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Math Skills Transparency Worksheets Chemistry: Matter and Change • Chapter 13 15

1. Describe how pressure relates to volume and temperature in the combined gas law.

2. How does volume relate to temperature in the combined gas law?

3. What is the equation for the combined gas law?

4. Kelvin temperature is used in the combined gas law, not Celsius. How are Celsiusdegrees converted to Kelvin degrees?

5. What is the first step in solving a combined gas law problem?

6. How is the equation for the combined gas law rearranged before solving it?

7. What is done to similar units in the numerator and the denominator in the final solving step?



Solving Gas Problems Using the Combined Gas LawSolving Gas Problems Using the Combined Gas Law



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Solving Gas Problems Using the Ideal Gas LawSolving Gas Problems Using the Ideal Gas Law



21

ProblemCalculate the number of moles of gas contained in a 2.0-L container at 200.0 K with a pressure of 120 kPa.


Known VariablesV = 2.0 L T = 200.0 KP = 120 kPa R = 8.314 L•kPa/(mol•K)Unknown Variablen = ? mol

Step 2. Solve for the unknown.Divide both sides of the ideal gas law equation by RT to solve for n.

PV � nRT n �

Substitute the known values into the rearranged equation.

n �(120 kPa)(2.0 L)

( )(200.0 K)

Multiply and divide numbers and units to solve for n.

n �(120 kPa)(2.0 L)

� 0.14 mol( )(200.0 K)


8.314�L�kPa��mol�K

8.314�L�kPa��mol�K

PV�RT

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1. What four variables does the ideal gas law describe?

2. What is the equation for the ideal gas law?

3. What is the first step in solving a combined gas law problem?

4. What numerical value of the gas constant, R, is used to solve this problem? Explain your answer.



7. If you were asked to find the molar mass of a gas in an ideal gas law problem, what formof the ideal gas law equation would you use? Show how this equation is derived from theoriginal form of the ideal gas law equation.

8. What is Avogadro’s principle and how does it relate to the ideal gas law?

Solving Gas Problems Using the Ideal Gas LawSolving Gas Problems Using the Ideal Gas Law



21

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18 Chemistry: Matter and Change • Chapter 13 Study Guide

GasesGases

Section 13.1 The Gas LawsIn your textbook, read about the basic concepts of the three gas laws.

Use each of the terms below to complete the passage. Each term may be used more than once.

Boyle’s law relates (1) and (2) if

(3) and amount of gas are held constant. Charles’s law relates

(4) and (5) if (6)

and amount of gas are held constant. Gay-Lussac’s law relates (7)

and (8) if (9) and amount of gas are

held constant.

In your textbook, read about the effects of changing conditions on a sample of gas.

For each question below, write increases, decreases, or stays the same.

10. The room temperature increases from 20°C to 24°C. Whathappens to the pressure inside a cylinder of oxygen contained inthe room?

11. What happens to the pressure of the gas in an inflated expandableballoon if the temperature is increased?

12. An aerosol can of air freshener is sprayed into a room. Whathappens to the pressure of the gas if its temperature staysconstant?

13. The volume of air in human lungs increases before it is exhaled.What happens to the temperature of the air in the lungs to causethis change, assuming pressure stays constant?

14. A leftover hamburger patty is sealed in a plastic bag and placed inthe refrigerator. What happens to the volume of the air in the bag?

15. What happens to the pressure of a gas in a lightbulb a few minutesafter the light is turned on?

STUDY GUIDE CHAPTER 13

pressure temperature volume

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Study Guide Chemistry: Matter and Change • Chapter 13 19Study Guide Chemistry: Matter and Change • Chapter 13 19

Section 13.2 The Combined Gas Law and Avogadro’s PrincipleIn your textbook, read about the combined gas law.

Fill in the following table. State what gas law is derived from the combined gas law when the variable listed in the first column stays constant and the variables in the second column change.

In your textbook, read about the relationships among temperature, pressure, and vol-ume of a sample of gas.

Fill in the blanks between the variables in the following concept map to show whetherthe variables are directly or inversely proportional to each other. Write direct or inversebetween the variables.

In your textbook, read about the combined gas law and Avogadro’s principle.

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

7. The variable that stays constant when using the combined gas law is

a. amount of gas. b. pressure. c. temperature. d. volume.

8. The equation for the combined gas law can be used instead of which of the followingequations?

a. Boyle’s law b. Charles’s law c. Gay-Lussac’s law d. all of these

9. Which of the following expresses Avogadro’s principle?

a. Equal volumes of gases at the same temperature and pressure contain equal numbers of particles.

b. One mole of any gas will occupy a certain volume at STP.

c. STP stands for standard temperature and pressure.

d. The molar volume of a gas is the volume that one mole occupies at STP.

temperature

volume pressure

5.

6.

4.

STUDY GUIDECHAPTER 13

Stays constant Change Becomes this law

Volume Temperature, pressure 1.

Temperature Pressure, volume 2.

Pressure Temperature, volume 3.

Derivations from the Combined Gas Law

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20 Chemistry: Matter and Change • Chapter 13 Study Guide 20 Chemistry: Matter and Change • Chapter 13 Study Guide

Answer the following questions.

10. What is standard temperature and pressure (STP)?

11. What is the molar volume of a gas equal to at STP?

In your textbook, read about how to solve problems using the combined gas law andAvogadro’s principle.

Each problem below needs more information to determine the answer. List as many letters as are needed to solve the problem.

a. molar volume of the gas d. pressure of the gas

b. molar mass of the gas e. volume of the gas

c. temperature of the gas f. No further information is needed.

12. What volume will 1.0 g N2 gas occupy at STP?

13. What volume will 2.4 mol He occupy at STP?

14. A gas sample occupies 3.7 L at 4.0 atm and 25°C. What volume will thesample occupy at 27°C?

15. A sample of carbon dioxide is at 273 K and 244 kPa. What will its volumebe at 400 kPa?

16. A sample of oxygen occupies 10.0 L at 4.00 atm pressure. At whattemperature will the pressure equal 3.00 atm if the final volume is 8.00 L?

17. At what pressure will a sample of gas occupy 5.0 L at 25°C if it occupies3.2 L at 1.3 atm pressure and 20°C?

18. How many grams of helium are in a 2-L balloon at STP?

19. One mole of hydrogen gas occupies 22.4 L. What volume will the sampleoccupy if the temperature is 290 K and the pressure is 2.0 atm?

Section 13.2 continued


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Section 13.2 The Ideal Gas LawIn your textbook, read about the ideal gas law.


1. Why is the mathematical relationship among the amount, volume, temperature, and pres-sure of a gas sample called the ideal gas law?

2. Define the ideal gas constant, R.

3. In Table 14.1 in your textbook, why does R have different numerical values?

4. What variable is considered in the ideal gas law that is not considered in the combinedgas law?

In your textbook, read about real versus ideal gases.

For each statement below, write true or false.

5. An ideal gas is one whose particles take up space.

6. At low temperatures, ideal gases liquefy.

7. In the real world, gases consisting of small molecules are the only gasesthat are truly ideal.

8. Most gases behave like ideal gases at many temperatures and pressures.

9. No intermolecular attractive forces exist in an ideal gas.

10. Nonpolar gas molecules behave more like ideal gases than do gasmolecules that are polar.

11. Real gases deviate most from ideal gas behavior at high pressures and lowtemperatures.

12. The smaller the gas molecule, the more the gas behaves like an ideal gas.


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In your textbook, read about applying the ideal gas law.

Rearrange the ideal gas law, PV � nRT, to solve for each of the following variables.Write your answers in the table.

In your textbook, read about using the ideal gas law to solve for molar mass, mass, or density.

Use the following terms below to complete the statements. Each term may be used more than once.

The number of moles of a gas is equal to the (17) divided by the

(18) .

Density is defined as (19) per unit (20) .

To solve for M in the equation M � , the (21) and the

(22) of the gas must be known.

According to the equation D � , the (23) of the gas must be

known when calculating density.

MP�RT

mRT�PV



Variable to Find Rearranged Ideal Gas Law Equation

n 13.

P 14.

T 15.

V 16.

Rearranging the Ideal Gas Law Equation

mass molar mass volume

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Section 13.3 Gas StoichiometryIn your textbook, read about gas stoichiometry.

Balance the following chemical equation. Then use the balanced equation to answer thequestions.

1. H2(g) � O2(g) 0 H2O(g)

2. List at least two types of information provided by the coefficients in the equation.

3. If 4.0 L of water vapor is produced, what volume of hydrogen reacted? What volume of oxygen?

4. If it is known that 2 mol of hydrogen reacts, what additional information would you needto know to find the volume of oxygen that would react with it?

5. List the steps you would use to find the mass of oxygen that would react with a knownnumber of moles of hydrogen.

6. Find the mass of water produced from 4.00 L H2 at STP if all of it reacts. Show your work.


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30 Chemistry: Matter and Change • Chapter 13

Student Recording Sheet

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.

Short Answer

Answer each question with complete sentences.

8.

9.

10.

Extended Response


11.

SAT Subject Test:Chemistry

12. 14.

13.

CHAPTER 13

Assessment

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Chapter 14 Mixtures and SolutionsMiniLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 32

ChemLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 33

Teaching Transparency Masters and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . . . 36

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

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



Table ofContents

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mini LAB 14Examine Freezing Point Depression

How do you measure freezing point depression?

Materials 400-mL beakers (2), crushed ice, rock salt (NaCl), cold tap water, stirring rods (2), nonmercury thermometer.

Procedure1. Read and complete the lab safety form..

2. Fill two 400-mL beakers with crushed ice. Add 50 mL of cold tap water to each beaker.

3. Measure the temperature of each beaker using a nonmercury thermometer.

4. Stir the contents of each beaker with a stirring rod until both beakers are at a constant temperature approximately 1 min. Record the temperature.

5. Add 75 g of rock salt (NaCl) to one of the beakers. Continue stirring both beakers. Some of the salt will dissolve.

6. When the temperature in each beaker is constant, record the final readings.

7. To clean up, flush the contents of each beaker down the drain with excess water.

Analysis

1. Compare your readings taken for the ice water and the salt water. How do youexplain the observed temperature change?

2. Explain why salt was added to only one of the beakers.

3. Explain Salt is a strong electrolyte that produces two ions, Na� and Cl�, when it dissociates in water. Explain why this is important to consider when calculating thecolligative property of freezing point depression?

4. Predict if it would be better to use coarse rock salt or fine table salt when makinghomemade ice cream. Explain.

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CHEMLAB 14

Safety Precautions

ProblemHow do factors affect the rate of solution formation?

Objectives• Observe the formation of

solutions• Hypothesize how stirring,

surface area, and temperature affect solution formation

• Infer the reasons that stirring, surface area, andtemperature affect solution formation

• Collect and analyze data

Materialscopper (II) sulfate

pentahydratedistilled water test tubes (6)25-mL graduated

cylinderglass stirring rod forceps test-tube rackmortar and pestlespatula clock

Investigate FactorsAffecting SolubilityBackground: The process of making a solution

involves the solvent coming in contact with thesolute particles. When you add a soluble compoundto water, several factors affect the rate of solutionformation.

Pre-Lab


2. How is forming an aqueous solution differentfrom a chemical reaction that takes place in aque-ous solution? Do you think that there are anysimilarities between the two processes? Explain.

3. What is the definition of temperature? When thetemperature of a solution goes up, what happensto the particles that make up the solution?

4. What is the surface area of a cube that is 8 cm ona side? Predict how the surface area will changeif the 8-cm-a-side cube is divided into 8 cubesthat are each 4 cm on a side. Calculate the newsurface area to verify your prediction.

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Procedure


2. Create a table to record your data.

3. Write a hypothesis that uses what you knowabout reaction rates to explain what you mightobserve during the procedure.

4. Place the six test tubes in the test-tube rack.

5. Place one crystal of copper (II) sulfate pentahy-drate in each of the first two test tubes.

6. For the remaining test tubes, use the mortar andpestle to crush a crystal. Use the spatula toscrape it into the third test tube.

7. Measure 15-mL of room-temperature distilledwater. Pour the water into the first test tube andrecord the time.

8. Observe the solution in the test tube just afteradding the water and after fifteen minutes.

9. Leave the first test tube undisturbed in the rack.

10. Repeat steps 7 and 8 for the third and fourthtest tubes.

11. Use the glass stirring rod to agitate the secondtest tube for 1-2 min.

12. Leave the third test tube undisturbed.

13. Agitate the fourth test tube with the glass stirring rod for 1-2 min.

14. Repeat steps 7 and 8 for the fifth test tube usingcold water. Leave the fifth test tube undis-turbed.

15. Repeat steps 7 and 8 for the sixth test tubeusing hot water. Leave the test tube undis-turbed.

16. Cleanup and Disposal Dispose of the remaining solids and solutions as directed byyour teacher. Wash and return all lab equipmentto its designated location.

CHEMLAB 14


1. Compare and Contrast What effect did you observe due to the agitation of the second and fourth testtubes verses the solutions in the first and third test tubes?

2. Observe and Infer What factors caused the more rapid solution formation in the fourth test tube incomparison to the second test tube?

3. Recognize Cause and Effect Why do you think the results for the third, fifth, and sixth test tubeswere different?

4. Discuss whether or not your data supported your hypothesis.

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5. Error Analysis Identify a major potential source of error for this lab and suggest an easymethod to correct it.

Inquiry Extension

Think Critically The observations made in this lab were macroscopic in nature. Propose asubmicroscopic explanation to account for these factors that affected the rate of solutionformation. At the molecular level, what is occuring to speed solution formation in each case?

CHEMLAB 14

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0 10 20 30 40 50 60 70 80 90 100

Solu

bili

ty (

g s

olu

te/1

00 g

H2O

) 90

100

80

70

60

50

40

30

20

10

0

Temperature (oC)

Solubilities as a Function of Temperature

NaClKClO3

KCl

CaCl2

Ce2(SO4)3

Solubility–Temperature GraphsSolubility–Temperature Graphs



42

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1. What variables are plotted on the graph?

2. What is the unit of each variable?

3. Use the graph to complete the table below.

4. At what temperature are sodium chloride and potassium chloride equally soluble

in water?

5. How does the solubility of cerium(III) sulfate differ from the solubility of potassium chlorate over the temperature range 0°C–100°C?

6. How many grams of sodium chloride will dissolve in 1.0 kg of water at 20°C?

7. Explain whether increasing temperature has a greater effect on the solubility of KCl oron the solubility of NaCl.

8. Explain how you might make a solution containing 42 g KCl dissolved in 100 g H2O at atemperature of 40°C. What term describes this type of solution?

Solubility–Temperature GraphsSolubility–Temperature Graphs



42

Substance Solubility at 10°C

Calcium chloride (CaCl2)

Cerium(III) sulfate (Ce2(SO4)3

Potassium chloride (KCl)

Potassium chlorate (KClO3)

Sodium chloride (NaCl)

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Phase Diagram of Solvent and SolutionPhase Diagram of Solvent and Solution



43

Pure solventSolution

LIQUIDSOLID

GAS

P

Tf Tb

Increasingtemperature

Freezing pointof solution

1 atm

Boiling pointof solution

Normal freezing point of water

Normal boiling point of water

Incr

easi

ng

p

ress

ure

� �

�

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1. What variables are plotted on the phase diagram?

2. What solvent is represented in the phase diagram?

3. What phases of the solvent are represented in the diagram?

4. What do the solid lines represent?

5. What is the term applied to a solution in which water is the solvent?

6. What do the dashed lines represent?

7. At each temperature, what does �P represent?

8. At any temperature, how does the vapor pressure of the aqueous solution compare withthe vapor pressure of the pure solvent?

9. Will a solution boil at the same temperature as the pure solvent under normal atmos-pheric pressure? Explain.

10. What must you do to the temperature of a solution to make it boil if it is at the boilingpoint of the pure solvent under normal atmospheric pressure?

11. How does the freezing point of a solution compare with the freezing point of the puresolvent at the same pressure?

Phase Diagram of Solvent and SolutionPhase Diagram of Solvent and Solution



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Percent by Mass

Percent by mass � � 100mass of solute

mass of solution

mass of solute � mass of solvent

Molality

Molality (m) �moles of solute

kilogram of solvent

Mole Fraction

XB �nB

nA � nB

mass of solvent (g)

molar mass of solvent (g/mol)

mass of solute (g)

molar mass of solute (g/mol)

1 kg

1000 gmass of solvent (g) �

mass of solute (g)

molar mass of solute (g/mol)

Calculating Percent by Mass,Mole Fraction, and Molalityfrom Mass Measurements




22

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1. Calculate the percent by mass, mole fraction, and molality of a solution that contains25.0 g of sodium chloride (NaCl) dissolved in 100.0 g of water.

Percent by Mass:

Mole Fraction:

Molality:





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Calculating the Boiling Point of a Solution

Tbsolution � Tbsolvent

� msoluteKbsolvent

Tbsolution � Tbsolvent

� msoluteKbsolvent

Tbsolution � 100°C � (2.00m)(0.515°C/m) � 100°C � 1.03°C

Tbsolution � 101.03°C

Find the boiling point of a 1.00m aqueous solution of KBr.

(Kbsolvent

� 0.515°C/m) Because a 1.00m aqueous solution of

potassium bromide contains 1 mol of K� ions and 1 mol of

Cl� ions per 1 kg of water, msolute � 2 � 1.00m, or 2.00m.

Boiling point of solution Boiling point elevation constantMolality of solute

�Tb: Boiling point elevationBoiling point of solvent

Calculating the Freezing Point of a Solution

Tfsolution � Tfsolvent

� msoluteKfsolvent

Tfsolution � Tfsolvent

� msoluteKfsolvent

Tfsolution � 0°C � (1.5m)(1.85°C/m) � 0°C � 2.78°C

Tfsolution � �2.78°C

Find the freezing point of a 1.50m aqueous solution of sucrose.

(Kfsolvent � 1.85°C/m) Because a 1.50m aqueous solution of sucrose

contains 1.50 mol of sucrose molecules per 1 kg of water, msolute � 1.50m.

Freezing point of solution Freezing point depression constant

Molality of solute particles

�Tf: Freezing point depressionFreezing point of solvent

Calculating Boiling Points andFreezing Points of SolutionsCalculating Boiling Points andFreezing Points of Solutions



23

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1. What is the boiling point of a 1.5m aqueous solution of fructose (C6H12O6)? The boilingpoint elevation constant of water is 0.515°C/m. Assume solute is a nonelectrolyte.

2. What is the freezing point of a solution of 1.5m aqueous solution of potassium bromide (KBr)?The freezing point depression constant of water is 1.86°C/m. Assume 100% dissociation.

3. Calculate the freezing point of a solution of a 0.975m aqueous solution of calcium nitrate(Ca(NO3)2). Assume 100% dissociation.

4. What is the boiling point of a solution of 20.0 g carbon tetrachloride (CCl4) dissolved in250.0 g of benzene (C6H6)? The boiling point of benzene is 80.10°C, and its boilingpoint elevation constant is 2.53°C/m. Assume solute is a nonelectrolyte.

5. Calculate the boiling point for a solution of 25.0 g calcium nitrate (Ca(NO3)2) dissolvedin 100.0 g of water. Assume 100% dissociation.

Calculating Boiling Points andFreezing Points of SolutionsCalculating Boiling Points andFreezing Points of Solutions



23

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Mixtures and SolutionsMixtures and Solutions


Section 14.1 Heterogeneous and Homogeneous MixturesIn your textbook, read about suspensions and colloids.


1. A solution is a mixture containing particles that settle out of the mixture ifleft undisturbed.

2. The most abundant substance in a colloid is the dispersion medium.

3. A colloid can be separated by filtration.

4. A solid emulsion consists of a liquid dispersed in a solid.

5. Whipped cream is an example of a foam.

6. In an aerosol, the dispersing medium is a liquid.

7. Brownian motion results from the collisions of particles of the dispersionmedium with the dispersed particles.

8. Dispersed particles in a colloid do not tend to settle out because they havepolar or charged atomic groups on their surfaces.

9. Stirring an electrolyte into a colloid stabilizes the colloid.

10. Colloids demonstrate the Tyndall effect.

The table below lists the characteristics of particles in colloids, solutions, and suspen-sions. Place a check in the column of each mixture whose particles have a particularcharacteristic.

Characteristics of Particles Colloid Solution Suspension

11. Less than 1 nm in diameter

12. Between 1 nm and 1000 nm in diameter

13. More than 1000 nm in diameter

14. Settle out if undisturbed

15. Pass through standard filter paper

16. Lower vapor pressure

17. Scatter light

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Study Guide Chemistry: Matter and Change • Chapter 14 45

Section 14.2 Solution ConcentrationIn your textbook, read about expressing concentration and using percent to describeconcentration.

Data related to aqueous solutions of sodium chloride (NaCl) and aqueous solutions ofethanol (C2H5OH) are provided in the table below. Use the table to answer the followingquestions. Circle the letter of the choice that best answers the question.

1. What is the percent by mass of NaCl in solution 1?

a. 0.030% b. 2.9% c. 3.0% d. 33%

2. Which of the following solutions is the most dilute?

a. Solution 1 b. Solution 2 c. Solution 3 d. Solution 4

3. What is the percent by volume of C2H5OH in Solution 5?

a. 0.2% b. 1.9% c. 2.0% d. 22%

4. Which of the following solutions is the most concentrated?

a. Solution 5 b. Solution 6 c. Solution 7 d. Solution 8

In your textbook, read about molarity and preparing molar solutions.

Read the following problem and then answer the questions.

An 85.0-mL aqueous solution contains 7.54 g iron(II) chloride (FeCl2). Calculate the molarityof the solution.

5. What is the mass of the solute?

6. What is the volume of the solution?

7. Write the equation that is used to calculate molarity.

8. In what unit must the amount of the solute be expressed to calculate molarity?

9. In what unit must the volume of the solution be expressed to calculate molarity?

10. Write the expression needed to convert the volume of the solution given in the problem

to the volume needed to calculate molarity.


Mass (g) Volume (mL)

Solution NaCl H2O Solution C2H5OH H2O

1 3.0 100.0 5 2.0 100.0

2 3.0 200.0 6 5.0 100.0

3 3.0 300.0 7 9.0 100.0

4 3.0 400.0 8 15.0 100.0

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11. What quantity must be used to convert the mass of the solute given in the problem to theamount of solute needed to calculate molarity?

12. Write the expression used to calculate the amount of solute.

13. Calculate the molarity of the solution. Show all your work.

In your textbook, read about molality and mole fractions.


14. How does molality differ from molarity?

15. Calculate the molality of a solution of 15.4 g sodium bromide (NaBr) dissolved in 125 gof water. Show all your work.

16. What is mole fraction?

17. Calculate the mole fraction of HCl in an aqueous solution that contains 33.6% HCl bymass. Show all your work.



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Section 14.3 Solvation and Solubility?In your textbook, read about the characteristics of solutions.

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

Air is a(n) (1) of oxygen gas dissolved in nitrogen

gas. The oxygen in air is the (2) , and nitrogen is the

(3) . Because oxygen gas dissolves in a solvent, oxygen gas

is a(n) (4) substance. A substance that does not dissolve is

(5) . (6) solutions are the most common

type of solutions. If one liquid is soluble in another liquid, such as acetic acid in water, the

two liquids are (7) . However, if one liquid is insoluble in another,

the liquids are (8) .

Read about solvation in aqueous solutions in your textbook.

The diagram shows the hydration of solid sodium chloride to form an aqueous solution.Use the diagram to answer the following questions.

9. Hydration is solvation in which the solvent is water. What is solvation?

Cl�

Cl�

Cl�

Cl� Cl�Na�

Na�

Na�

Na�

O HH

Na�

Cl�

Cl�

� �

�

immiscible liquid soluble solution

insoluble miscible solute solvent

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10. As sodium chloride dissolves in water, what happens to the sodium and chloride ions?

11. Explain the orientation of the water molecules around the sodium ions and chloride ions.

12. How does the strength of the attraction between water molecules and sodium and chloride ions compare with the strength of the attraction between the sodium ions and chloride ions? How do you know?

13. List three ways that the rate of solvation may be increased.

In your textbook, read about heat of solution, solubility, and factors that affect solubility.


14. The overall energy change that occurs when a solution forms is called theheat of solution.

15. Solubility is a measure of the minimum amount of solute that dissolves ina given amount of solvent at a specified temperature and pressure.

16. Solvation continues as long as the solvation rate is less than thecrystallization rate.

17. In a saturated solution, solvation and crystallization are in equilibrium.

18. Additional solute can be dissolved in an unsaturated solution.

19. The solubility of a gas dissolved in a liquid decreases as the temperatureof the solution increases.



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Section 14.4 Colligative Properties of SolutionsIn your textbook, read about electrolytes and colligative properties, vapor pressure low-ering, boiling point elevation, and freezing point depression.

Use the table to answer the following questions.

1. Which properties in the table are colligative properties?

2. What can you conclude about the relationship between colligative properties and thenumber of ions in solution from the 1.0m NaCl(aq) and 2.0m NaCl(aq) solutions?

3. What can you conclude about the relationship between colligative properties and the typeof ions in solution from the 1.0m HCl(aq) and 1.0m NaCl(aq) solutions?

Suppose that in a simple system, a semipermeable membrane is used to separate asucrose-water solution from its pure solvent, water. Match the descriptions of the systemin Column A with the terms in Column B.

Column A Column B

4. Cannot cross the semipermeable membrane

5. Can cross the semipermeable membrane

6. The side that exerts osmotic pressure

7. The diffusion of the solvent particles across thesemipermeable membrane from the area of higher solventconcentration to the area of lower solvent concentration

8. The barrier with tiny pores that allow some particles topass through but not others

9. The side from which more water molecules cross thesemipermeable membrane

10. A colligative property of solutions

Solution Density (g/L) Boiling Point (°C) Freezing Point (°C)

1.0m C2H5OH(aq) 1.05 100.5 �1.8

1.0m HCl(aq) 1.03 101.0 �3.7

1.0m NaCl(aq) 1.06 101.0 �3.7

2.0m NaCl(aq) 1.12 102.1 �7.4

a. osmotic pressure

b. water molecules

c. semipermeable membrane

d. sugar molecules

e. osmosis

f. solution side

g. pure solvent side

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CHAPTER 14

Assessment



1. 4. 7. 10.

2. 5. 8.

3. 6. 9.

Short Answer


11.

12.

13.

Extended Response


14.

15.


16. 18.

17. 19.

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Chapter 15 Energy and Chemical ChangeMiniLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 58

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


Math Skills Transparency Masters and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . . . 68

Study Guide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 74



Table ofContents

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mini LAB 15Determine Specific Heat

How can you determine the specific heat of a metal? You can use a coffee-cupcalorimeter to determine the specific heat of a metal.

Materials distilled water, 250-mL beaker (2), hot plate, balance, metal cylinder, crucibletongs, graduated cylinder, polystyrene coffee cup, nonmercury thermometer


2. Make a table to record your data.

3. Pour approximately 150 mL of distilled water into a 250-mL beaker. Place the beaker ona hot plate set on high.

4. Use a balance to find the mass of a metal cylinder.

5. Using crucible tongs, carefully place the metal cylinder in the beaker on the hot plate.

6. Measure 90.0 mL of distilled water using a graduated cylinder.

7. Pour the water into a polystyrene coffee cup nested in a second 250-mL beaker.

8. Measure and record the temperature of the water using a nonmercury thermometer.

9. When the water on the hot plate begins to boil, measure and record the temperature asthe initial temperature of the metal.

10. Carefully add the hot metal to the cool water in the coffee cup with the crucible tongs.Do not touch the hot metal with your hands.

11. Stir, and measure the maximum temperature of the water after the metal was added.

Analysis

1. Calculate the heat gained by the water. The specific heat of H2O is 4.184 J/g·°C.Because the density of water is 1.0 g/mL, use the volume of water as the mass.

2. Calculate the specific heat of your metal. Assume that the heat absorbed by thewater equals the heat lost by the metal.

3. Compare this experimental value to the accepted value for your metal.

4. Describe major sources of error in this lab. What modifications could you make inthis experiment to reduce the error?

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CHEMLAB 15

Safety Precautions• Always wear safety goggles and a lab apron.• Tie back long hair.• Hot objects may not appear to be hot.• Do not heat broken, chipped, or cracked glassware.• Do not eat any items used in the lab.

ProblemHow many Calories are in a potato chip?

Objectives• Identify the reactants and

products in the reaction.• Measure mass and temper-

ature in order to calculatethe amount of heatreleased in the reaction.

• Propose changes in theprocedure and design ofthe equipment todecrease the percenterror.

Measure CaloriesThe burning of a potato chip releases heat stored in the

substances contained in the chip. Using calorimetry, youwill approximate the amount of energy contained in a potatochip.

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. Form a hypothesis about how the quantity of heatproduced by the combustion reaction will com-pare with the quantity of heat absorbed by thewater.

4. What formula will you use to calculate the quan-tity of heat absorbed by the water?

5. Assuming that the potato chip contains compoundsmade up of carbon and hydrogen, what gases willbe produced in the combustion reaction?

Materialslarge potato chipor other snack food250-mL beaker 100-mL graduated

cylinder evaporating dishnonmercury

thermometerring stand with ringwire gauzematches

stirring rodbalance

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Procedure


2. Measure the mass of a potato chip and record itin a data table.

3. Place the potato chip in an evaporating dish onthe metal base of the ring stand. Position the ringand wire gauze so that they will be 10 cm abovethe top of the potato chip.

4. Measure the mass of an empty 250-mL beakerand record it in the data table.

5. Using a graduated cylinder, measure 50 mL ofwater and pour it into the beaker. Measure themass of the beaker and water and record it inyour data table.

6. Measure and record the initial temperature of thewater.

7. Place the beaker on the wire gauze on the ring stand. Use a match to ignite the bottom ofthe potato chip.

8. Gently stir the water in the beaker while the chipburns. Measure and record the highest tempera-ture attained by the water.

9. Cleanup and Disposal Wash all lab equipmentand return it to its designated place.

CHEMLAB 15


1. Classify Is the reaction exothermic or endothermic? Explain how you know.

2. Observe and Infer Describe the reactant and products of the chemical reaction. Was the reactant(potato chip) completely consumed? What evidence supports your answer?

3. Calculate Determine the mass of the water and its temperature change. Use the equation q � c � m � �T to calculate how much heat, in joules, was transferred to the water in thebeaker by the burning of one chip.

Mass of beaker and 50 mL of water

Mass of empty beaker

Mass of water in beaker

Mass of potato chip

Highest temperature of water

Initial temperature of water

Change in temperature

Observations of the Burning of a Potato Chip

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4. Calculate Convert the quantity of heat in joules/chip to Calories/chip.

5. Calculate From the information on the chip’s container, determine the mass in grams ofone serving. Determine how many Calories are contained in one serving. Use your data tocalculate the number of Calories released by the combustion of one serving.

6. Error Analysis Compare your calculated Calories per serving with the value on thechip’s container. Calculate the percent error.

7. Compare your class results with other sutdents by posting your data at glencoe.com

Inquiry Extension

Predict Do all potato chips have the same number of calories? Make a plan to test severaldifferent brands of potato chips.

CHEMLAB 15

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Thermometer

Stirrer

Ignitionterminals

Insulation

Sealedreactionchambercontainingsubstanceand oxygen

Water

Using a CalorimeterUsing a Calorimeter



44

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1. The calorimeter shown on the transparency is used to measure the caloric content offoods. To do this, a sample of food is burned inside the reaction chamber of the calorime-ter. What is the system? What are the surroundings?

2. What besides food must be added to the chamber? Explain why.

3. What are the products of the reaction that takes place in the reaction chamber?

4. Why is the calorimeter insulated?

5. What does the thermometer measure?

6. Describe the movement of heat as the reaction takes place inside the chamber.

7. Assuming that no heat escapes from the calorimeter, what equation would you use todetermine the amount of heat released by the burning food in the reaction chamber?Define all variables in the equation.

8. Does the answer obtained from the equation in question 7 have a positive or negativevalue? Explain why. What is the sign of �H for the reaction?

Using a CalorimeterUsing a Calorimeter



44

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120

100 80 60 40 20

2010

020

0H

eat

Ad

ded

(kc

al)

Liq

uid

wat

er a

nd

wat

er v

apo

r

Liq

uid

wat

erIc

e an

dliq

uid

wat

er

Ice

Water vapor

740

Temperature (�C)

0 0

–20

–40

Temperature Changes of WaterTemperature Changes of Water



45

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1. The graph shows what happens when ice, at �40°C, is gradually heated to more than120°C. What is happening in the region between �40°C and 0°C?

2. At 0°C, the temperature does not change even though 80 kcal of heat is added. Why doesthe temperature not change?

3. What is happening in the region between 0°C and 100°C?

4. When the water reaches 100°C, the temperature does not change even though 540 kcal ofheat is added. Why does the temperature not change?

5. Compare the amount of heat needed to convert liquid water to water vapor with theamount needed to convert ice to liquid water. Explain the difference.

6. What is happening in the region above 100°C?

7. If you continue to add heat, what will happen to the water vapor?

Temperature Changes of WaterTemperature Changes of Water



45

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66 Chemistry: Matter and Change • Chapter 15 Teaching Transparency Master

H2(

g) �

1 O

2(g

)2

�H

� �

242

kJ

�H

� �

286

kJ

�H

� �

44 k

J

H2O

(g)

H2O

(l)

Changes in Enthalpy and EntropyChanges in Enthalpy and Entropy


Use with Chapter 15,Sections 15.4 and 15.5

46

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1. What do the arrows on the transparency represent?

2. Why do the arrows vary in direction?

3. What can you conclude about the transitions and the magnitudes of the enthalpies shownon the transparency?

4. How does Hess’s law apply to your answer to Question 3?

5. What do the ball models for liquid water and gaseous water on the transparency show?

6. What do the ball models indicate about the overall order of the molecules?

7. When molecules become more ordered or disordered, what happens to the entropy?

Changes in Enthalpy and EntropyChanges in Enthalpy and Entropy


Use with Chapter 15,Sections 15.4 and 15.5

46

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Determining the Heat Capacityof a CalorimeterDetermining the Heat Capacityof a Calorimeter



24

• Place 100.0 g of room temperature water in the calorimeter. Let thewater and calorimeter come to a constant temperature. Record thistemperature as: Ti cold water.

• Then add 100.0 g of hot water to the calorimeter. Record thetemperature of the hot water as: Ti hot water.

• Measure the temperature in the calorimeter every 15 seconds for 10 minutes.

• Plot temperature versus time. From the graph, determine the finaltemperature, Tf.

To determine the heat capacity of the calorimeter, use therelationship:

�Heathot water � �Heatcold water � Heatcalorimeter

�(masshot water � (specific heat of water) � (Tf � Ti hot water)) �

�masscold water � (specific heat of water) � (Tf � Ti cold water )

� Heat Capacity � (Tf � Ti cold water)

Example:

�(100.0 g � 4.184 J/g�°C � (29.8°C � 40.0°C)) �

(100.0 g � 4.184 J/g�°C � (29.8°C � 20.0°C)) � (Heat Capacity � (29.8°C � 20.0°C))

4267.68 J � 4100.32 J � (Heat Capacity � (9.8°C))

167.36 J / 9.8°C � Heat Capacity

17.1 J/°C � Heat Capacity

A calorimeter relies on the fact that the amount of heat lost by ahot body equals the amount of heat gained by a cold body. Theamount of heat absorbed or lost by the calorimeter apparatus isrelated to its heat capacity and the temperature change.

Mass Initial Temperature Final Temperature

Hot water 100.0 g 40.0°C 29.8°C

Cold water 100.0 g 20.0°C 29.8°C

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1. Why is it unnecessary to know the mass of the calorimeter apparatus to determine itsheat capacity?

2. You are using a calorimetry experiment to determine the amount of heat absorbed when asolid melts. How would the results be affected if the heat capacity of the calorimeter wasnot included in your calculations? Why?

3. What is the mathematical relationship for heat from the hot water?

4. What is the mathematical relationship for heat from the cold water?

5. Why does the calorimeter apparatus use a heat capacity rather than a specific heat?

6. Determine the heat capacity of a calorimeter using the following data.

Determining the Heat Capacityof a CalorimeterDetermining the Heat Capacityof a Calorimeter



24

Mass Initial Temperature Final Temperature

Hot water 150.0 g 45.0°C 36.8°C

Cold water 100.0 g 25.0°C 36.8°C

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Hess’s LawHess’s Law



25

How can you find the enthalpy of a specificreaction from a series of other reactions?

Here is a series of chemical reactions and their enthalpies.

H2(g) � Cl2(g) 0 2HCl(g) �H° � �185 kJ

2H2(g) � O2(g) 0 2H2O(g) �H° � �483.7 kJ

Calculate �H° for the reaction: 4HCl(g) � O2(g) 0 2Cl2(g) � 2H2O(g)

Reverse the reaction that contains HCl(g) to make HCl(g) a reactant.Change the sign of its enthalpy value because the reaction is reversed.

2HCl(g) 0 H2(g) � Cl2(g) �H° � 185 kJ

Multiply the reversed equation and enthalpy by 2 to obtain 4 mol HCl reacting.

2(2HCl(g) 0 H2(g) � Cl2(g) �H° � �185 kJ)

Reversing the equation and the sign of the enthalpy gives equation 3.

4HCl(g) 0 2H2(g) � 2Cl2(g) �H° � �370 kJ

Reaction 2 already has O2(g) as a reactant and the coefficients have thecorrect values. The following two equations and their enthalpies can beadded to obtain the desired equation.

4HCl(g) 0 2H2(g) � 2Cl2(g) �H° � �370 kJ

2H2(g) � O2(g) 0 2H2O(g) �H° � �483.7 kJ

4HCl(g) � 2H2(g) � O2(g) 0 2H2(g) � 2Cl2(g) � 2H2O(g) �H° � �370 kJ �(�483.7 kJ)

The term 2H2(g) appears on both sides of the equation so it can be cancelled. Follow the rules for significant digits for the final enthalpyvalue.

4HCl(g) � O2(g) 0 2Cl2(g) � 2H2O(g) �H° � �114 kJ

2

3

3

2

1

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1. In the example, what type of reactions are described in the starting equations?

2. What must be done to the formation reaction for hydrogen chloride to obtain the targetequation?

3. What must be done to the formation reaction for water to obtain the target equation?

4. How can the enthalpy for the target reaction be obtained?

5. Calculate �H° for the target reaction: 2Al(s) � Fe2O3(s) 0 2Fe(s) � Al2O3(s). Showyour work. Use the following reactions and enthalpies.

4Al(s) � 3O2(g) 0 2Al2O3(s) �H° � �3351.4 kJ

4Fe(s) � 3O2(g) 0 2Fe2O3(s) �H° � �1648.4 kJ

Hess’s LawHess’s Law



25

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Determining SpontaneityDetermining Spontaneity



26

Determine whether a chemical reaction isspontaneous or nonspontaneous.

From the relationship, �G � �H � T�S, it can be seen that when �H isnegative and �S is positive the reaction will always be spontaneous.When �H is positive and �S is negative, the reaction will always benonspontaneous.

What happens when both �H and �S are positive or when they are bothnegative? In these cases, spontaneity depends on the temperature.

The condensation of a gas identified as A has an enthalpy of �2.00 kJ/mol (�2000 J/mol) and an entropy of �12.0 J/mol�K. When will this condensation be spontaneous and when will it be nonspontaneous?

Use the relationship �G � �H � T�S to determine spontaneity. Find thetemperature at which equilibrium exists, that is, when �G � 0. At lower temperatures, the condensation is spontaneous. At highertemperatures, the condensation is nonspontaneous.

At equilibrium, �G � 0. Therefore, �H � T�S.

Solving for T gives: T �

Substituting the values for the condensation of A gives: T �

T � 167 K

To be spontaneous, �G must be negative, and to be nonspontaneous, �Gmust be positive.

To check this, perform test calculations.

�G � �2.00 kJ/mol � � (166 K � (�12.0 J/mol�K)) � �8 J/mol

(spontaneous below 167 K)

�G � �2.00 kJ/mol � � (168 K � (�12.0 J/mol�K)) � �16 J/mol

(nonspontaneous above 167 K)

1000 J1 kJ

1000 J1 kJ

(�2.00 kJ/mol)(1000 J)(�12.0 J/mol�K)(1 kJ)

�H��S

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1. The term T�S always has the same sign as �S. Explain why this is so.

2. A gas has a very large �H and a very small �S. Both values have negative signs. Whatcan you say about the temperature below which condensation will be spontaneous?

3. A chemical reaction has a �H of �135.5 kJ/mol and a �S of �148.9 J/ mol�K. Will thisreaction be spontaneous at a temperature of 750°C? Show your work.

4. Use the data given in question 3 to determine the temperature at which the reaction willbe in equilibrium. Give your answer in kelvins. Show your work.

Determining SpontaneityDetermining Spontaneity



26

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74 Chemistry: Matter and Change • Chapter 15 Study Guide74 Chemistry: Matter and Change • Chapter 15 Study Guide

Energy and Chemical ChangeEnergy and Chemical Change

Section 15.1 EnergyIn your textbook, read about the nature of energy.

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. Energy is the ability to do work or produce heat.

2. The law of conservation of energy states that energy can becreated and destroyed.

3. Chemical potential energy is energy stored in a substance becauseof its composition.

4. Heat is a form of energy that flows from a warmer object to acooler object.

5. A calorie is the amount of energy required to raise the temperatureof one gram of pure water by one degree Celsius.

6. A calorie is the SI unit of heat and energy.

7. The specific heat of a substance is the amount of heat required toraise the temperature of one gram of that substance by one degreeCelsius.

8. Kinetic energy is energy of motion.

9. Chemicals participating in a chemical reaction contain only potential energy.

10. One nutritional Calorie is equal to 100 calories.

11. One calorie equals 4.184 joules.

12. When a fuel is burned, some of its chemical potential energy islost as heat.

13. To convert kilojoules to joules, divide the number of kilojoules by1000 joules/1 kilojoule.

Answer the following question. Show all your work.

14. If the temperature of a 500.0-g sample of liquid water is raised 2.00°C, how much heat isabsorbed by the water? The specific heat of liquid water is 4.184 J/(g�°C).


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a. system

b. calorimeter

c. thermochemistry

d. universe

e. enthalpy

f. enthalpy (heat) ofreaction

g. surroundings


Section 15.2 Heat in Chemical Reactions and ProcessesIn your textbook, read about measuring heat and about chemical energy and the universe.

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

Column A Column B

1. An insulated device used to measure the amount of heatabsorbed or released during a chemical or physicalprocess

2. The study of heat changes that accompany chemicalreactions and phase changes

3. The specific part of the universe that contains the reactionor process you wish to study

4. The change in enthalpy in a chemical reaction

5. A system plus its surroundings

6. The heat content of a system at constant pressure

7. Everything in the universe except the system beingstudied

Use the illustration to answer the following questions.

8. A scientist is studying the solution in the flask. What is the system?

9. What are the surroundings?

10. What is the universe?

Solution ofBa(OH)2

andNH4NO3

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Section 15.3 Thermochemical EquationsIn your textbook, read about writing thermochemical equations and about changes of state.

Use the following terms to complete the statements. Some terms will be used more than once.

1. A(n) is a balanced chemical equation thatincludes the physical states of all reactants and products and the energy changethat accompanies the reaction.

2. The enthalpy change for the complete burning of one mole of a substance is the

.

3. The is the heat required to vaporize one moleof a liquid.

4. The is the heat required to melt one mole of asolid substance.

5. Converting two moles of a liquid to a solid requires an amount of energy that is twice

the .

6. 2H2(g) � O2(g) 0 2H2O(g) �H � �572 kJ is a(n) .

7. The conversion of a gas to a liquid involves the

8. When a gas condenses to a liquid, heat is to thesurroundings.

9. Sweating makes you feel cooler because, as it evaporates, the water on your skin

heat from your body.

10. If you put an ice cube in a glass of soda pop, the heat absorbed by the ice will cause the

ice to melt, and the soda pop will become .

11. If it takes 100 joules to melt a piece of ice, must be absorbed by the ice.

12. In the equation H2O(s) 0 H2O(l) �H � 600 kJ, the positive value for �H means that

is absorbed in the reaction.


thermochemical equation enthalpy of combustion released

molar enthalpy of vaporization molar enthalpy of fusion absorbs

cool heat

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Section 15.4 Calculating Enthalpy ChangeIn your textbook, read about Hess’s law and standard enthalpy (heat) of formation.


1. Hess’s law states that if two or more thermochemical equationscan be added to produce a final equation for a reaction, then thesum of all the enthalpy changes for the individual reactions is theenthalpy change for the final reaction.

2. The standard enthalpy of formation is the change in enthalpy thataccompanies the formation of one gram of a compound in itsstandard state from its constituent elements in their standardstates.

3. The standard state of iron is solid.

4. For a pure gas, the standard state is the gas at a pressure of oneatmosphere.

5. The symbol used to represent standard enthalpy of formation is�Hf°.

6. The standard state of a substance is the normal state of thesubstance at 0 K and one atmosphere pressure.

7. The standard enthalpy of formation of a free element in itsstandard state is 0.0 kJ.

8. A standard enthalpy of formation that has a negative value meansthat energy is absorbed during the reaction.

9. The standard state of oxygen is gas.

10. Standard enthalpies of formation provide data for calculating theenthalpies of reactions under standard conditions using Hess’slaw.

11. The standard state of mercury is solid.


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Use the table to answer the following questions.

12. What does a formation equation show?

13. What does the negative sign on the value of an enthalpy of formation indicate?

14. Using the formation equations for CH4(g), CH3OH(g), and H2O(g), calculate �Hrxn forthe following equation. Show and explain all your work.

CH4(g) � H2O(g) 0 CH3OH(g) � H2(g)


Compound Formation Equation �Hf° (kJ/mol)

CH4(g) C(graphite) � 2H2(g) 0 CH4 (g) 75

CH3OH(g) C(graphite) � 2H2(g) � O2(g) 0 CH3OH(g) 239

H2O(g) O2(g) � H2(g) 0 H2O(g) 2421�2

1�2


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Section 15.5 Reaction SpontaneityIn your textbook, read about spontaneous processes and about entropy, the universe,and free energy.

Use each of the terms below to complete the statements.

1. A(n) is a physical or chemical change that occurs with no

outside intervention.

2. A measure of disorder or randomness of the particles that make up a system is called

.

3. The states that spontaneous processes always proceed in such

a way that the entropy of the universe increases.

4. is the energy that is available to do work.


5. A process cannot be spontaneous if it is exothermic and there is anincrease in disorder.

6. A process cannot be spontaneous if it is endothermic and there is adecrease in disorder.

7. A process cannot be spontaneous if it is exothermic and there is a decreasein disorder as long as the temperature remains low.

8. A process cannot be spontaneous if it is endothermic and there is anincrease in disorder as long as the temperature remains high.

9. A process can never be spontaneous if the entropy of the universeincreases.

10. When �G for a reaction is negative, the reaction is spontaneous.

11. When �G for a reaction is positive, the reaction is not spontaneous.

12. When �H for a reaction is negative, the reaction is never spontaneous.

13. When �H for a reaction is large and positive, the reaction is not expectedto be spontaneous.


spontaneous process entropy second law of thermodynamics free energy

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1. 4. 7.

2. 5. 8.

3. 6. 9.

Short Answer


10.

11.

12.

Extended Response


13.

14.


15. 17.

16. 18.

CHAPTER 15

Assessment

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Chapter 16 Reaction RatesMiniLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 88

ChemLab . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 89


Math Skills Transparency Masters and Worksheets . . . . . . . . . . . . . . . . . . . . . . . . . . 98

Study Guide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 100

Chapter Assessment. . . . . . . . . . . . . . . . . . . . . . . . . . . . 106

STP Recording Sheet . . . . . . . . . . . . . . . . . . . . . . . . . . . 112

Table ofContents

87

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mini LAB 16Examine Reaction Rate and Temperature

What is the effect of temperature on a common chemical reaction?

Materials: nonmercury thermometer, hot plate, 250-mL beaker, balance, water, effervescent tablet, stopwatch or clock with second hand


2. Break a single effervescent tablet into four equal pieces.

3. Use a balance to measure the mass of one piece of the tablet. Measure 50 mL of room-temperature water (approximately 20°C) into a 250-mL beaker. Use a nonmercurythermometer to measure the temperature of the water.

4. With a stopwatch or a clock with seond hand ready, add the piece of tablet to the water.Record the amount of time elapsed between when the tablet hits the water and when allof the solid has dissolved.

5. Repeat Steps 3 and 4, this time gradually warming the 50 mL of water to about 50°C on ahot plate. Maintain the temperature (equilibrate) throughout the run.

Analysis

1. Identify the initial mass the final mass, and t1 and t2 for each run.

2. Calculate the reaction rate by finding the mass of reactant consumed per second for each run.

3. Describe the relationship between reaction rate and temperature for this reaction.

4. Predict what the reaction rate would be if the reaction rate were carried out at 40°Cand explain the basis for your prediction. To test your prediction, repeat the reactionat 40°C using another piece of tablet.

5. Evaluate how well your prediction for the reaction rate at 40°C compares to themeasured reaction rate.

Mass of Reaction Reaction Temperature Tablet (g) time (s) rate (g/s) (°C)

Data Table

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CHEMLAB 16

Safety Precautions• Always wear safety goggles and a lab apron.• Never pipette any chemical by mouth.• Hydrochloric acid is corrosive. Avoid contact with skin and eyes.

ProblemHow does the concentra-tion of a reactant affectthe reaction rate?

Objectives• Sequence the acid con-

centrations from the mostto the least concentrated.

• Observe which concentra-tion results in the fastestreaction rate.

Materials10-mL graduated

pipette safety pipette filler6M hydrochloric

aciddistilled water25 mm � 150 mm

test tubes (4)test-tube rack

magnesium ribbonemery cloth or fine

sandpaperscissorsplastic rulertongswatch with second

hand or stop-watch

stirring rod

Observe How ConcentrationAffects Reaction RateThe collision theory describes how the change in concentration of

one reactant affects the rate of chemical reactions.

Pre-Lab

1. Read the entire CHEMLAB. Prepare all writtenmaterials that you will take into the laboratory.Be sure to include safety precautions and proce-dure notes. Use the data table on the next page.

2. Use emery paper or sandpaper to polish the mag-nesium ribbon until it is shiny. Use scissors to cutthe magnesium into four 1-cm pieces.

3. Place the four test tubes in the test-tube rack.Label the test tubes #1 (6.0M HCl), #2 (3M HCl), #3 (1.5M HCl), and #4 (0.75M HCl).

4. Form a hypothesis about how the chemical reac-tion rate is related to reactant concentration.

5. What reactant quantity is held constant? What arethe independent and dependent variables?

6. What gas is produced in the reaction betweenmagnesium and hydrochloric acid? Write the balanced formula equation for the reaction.

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7. Why is it important to clean the magnesium rib-bon? If one of the four pieces is not thoroughlypolished, how will the rate of the reaction involv-ing that piece be affected?

Procedure


2. Use a safety pipette filler to draw 10 mL of 6.0Mhydrochloric acid (HCl) into a 10-mL graduatedpipette.

3. Dispense the 10 mL of 6.0M HCl into Test Tube #1.

4. Draw 5.0 mL of the 6.0M HCl from Test Tube #1with the empty pipette. Dispense this acid intoTest Tube #2 and use the pipette to add an addi-tional 5.0 mL of distilled water to the acid. Usethe stirring rod to mix thoroughly. This solution is 3.0M HCl.

5. Draw 5.0 mL of the 3.0M HCl from Test Tube #2with the empty pipette. Dispense this acid intoTest Tube #3 and use the pipette to add an addi-tional 5.0 mL of distilled water to the acid. Usethe stirring rod to mix thoroughly. This solution is 1.5M HCl.

6. Draw 5.0 mL of the 1.5M HCl from Test Tube #3with the empty pipette. Dispense this acid intoTest Tube #4 and use the pipette to add an addi-tional 5.0 mL of distilled water to the acid. Usethe stirring rod to mix thoroughly. This solutionis 0.75M HCl.

7. Draw 5.0 mL of the 0.75M HCl from Test Tube#4 with the empty pipette. Neutralize and discardit in the sink.

8. Using the tongs, place a 1-cm length of magne-sium ribbon into Test Tube #1. Record the time inseconds that it takes for the bubbling to stop.

9. Repeat step 8 using the remaining three TestTubes of HCl and the three remaining pieces ofmagnesium ribbon. Record in the data table thetime (in seconds) it takes for the bubbling to stop.

9. Cleanup and Disposal Place acid solutions inan acid discard container. Thoroughly wash alltest tubes and lab equipment. Discard other mate-rials as directed by your teacher. Return all labequipment to its proper place.

CHEMLAB 16

Test tube [HCl] (M) Time (s)

1

2

3

4

Reaction Time Data

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1. Make a Graph Plot the concentration of the acid on the x-axis and the reaction time on the y-axis. Draw a smooth curve through the data points.

2. Conclude Based on your graph, what is the relationship between the acid concentration and the reaction rate?

3. Hypothesis Write a hypothesis using collision theory, reaction rate, and reactantconcentration to explain your results.

4. Error Analysis Compare your experimental results with those of several other students inthe laboratory. Explain the differences.

Inquiry Extension

Design an Experiment Based on your observations and results, would temperature variationsaffect reaction rates. Plan an experiment to test your hypothesis.

CHEMLAB 16

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Factors That Affect Reaction RateFactors That Affect Reaction Rate



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1. Look at the sequence of pictures. What happened to the apple over time?

2. No chemical was added to the apple. Explain why the apple changed.

3. How could you test your answer to Question 3?

4. What effect would increasing the amount of air surrounding the apple have on the apple?Explain your answer.

5. How would slicing the apple into more pieces affect the apple? Explain your answer.

6. The apple in the pictures is raw. A cooked apple would not change the same way. Give apossible reason why.

Factors That Affect Reaction RateFactors That Affect Reaction Rate



47

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Reaction OrderReaction Order



48

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1. What is formed when the two solutions mix?

2. Explain how the substance is formed in terms of the particles in the two solutions.

3. On the basis of the image in the transparency, what can you conclude about the rate ofthe reaction? Explain your answer.

4. On the basis of your answer to Question 3, what can you conclude about the activationenergy for the reaction?

5. What can you conclude about the reaction order for the reaction?

6. If the rate equation is determined to be k[Pb2�][I�], what is the reaction order?

7. Would you classify the reaction as exothermic, endothermic, nonspontaneous, or sponta-neous. Explain your answer.

Reaction OrderReaction Order



48

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1.00

0.80

0.60

[H2O

2] (

mo

l/L)

�[H

2O2]

0.40

0.20

00 1 2 3 4 5 6 7 8 9 10

Relative time

Change in [H2O2] with Time

�t

Reaction RateReaction Rate



49

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1. What does the graph show?

2. Is the hydrogen peroxide a reactant or product in the reaction? How do you know?

3. The hydrogen peroxide is part of a decomposition reaction. Write a balanced equation forthe reaction.

4. From the graph, what can you conclude about the rate of this reaction?

5. Define the term instantaneous rate.

6. How can you use the graph to find the instantaneous rate of the reaction that you identi-fied in your answer to Question 3? Give your answer in mathematical terms.

7. To show the complete relationship of reactants to products in this reaction, what elsewould you need to plot on the graph?

8. What would you expect your answer to Question 7 to show?

Reaction RateReaction Rate



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Determining Reaction Orders Determining Reaction Orders



27

Trial Initial [A] (M) Initial [B] (M) Initial Rate (mol/L�s)

1 0.30 0.10 1.20 � 10�2

2 0.60 0.10 4.80 � 10�2

3 0.60 0.20 9.60 � 10�2

Step 1 Write the rate equations for trials 1

and 2.

Step 2 Compare Rate 2 with Rate 1. (Hint:

Use the initial rate data.)

Step 3 Substitute the rate equations in

step 1 into the equation in step 2.

Step 4 Compare the concentrations of each

reactant. (Hint: Use the concentration

data.)

Step 5 Substitute the equations from step 4

into the equation of step 3 and

simplify.

Rate 1 � k[A1]m[B1]n

Rate 2 � k[A2]m[B2]n

Rate 2 � 4 Rate 1

k[A2]m[B2]n � 4k[A1]m[B1]n

[A2] � 2[A1] [B2] � [B1]

k(2[A1])m[B1]n � 4k[A1]m[B1]n

k(2)m[A1]m[B1]n � 4k[A1]m[B1]n

(2)m � 4

m � 2

What is the order of the reaction for reactant A?

Assume that the general rate law for this type of reaction is

Rate � k[A]m[B]n

Determine m by comparing trials 1 and 2.

Experimental Initial Rates for aA � bB 0 products

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Use the following data to determine the order of the reaction for each reactant.

Assume that the general rate law for this type of reaction is Rate � k[A]m[B]n

1. Determine n by comparing trials 1 and 2. Follow Steps 1–5 shown on the transparency.

Step 1.

Step 2.

Step 3.

Step 4.

Step 5.

2. Determine m by comparing trials 2 and 3. Follow Steps 1–5 shown on the transparency.

Step 1.

Step 2.

Step 3.

Step 4.

Step 5.

Determining Reaction Orders Determining Reaction Orders



27

Experimental Initial Rates for aA � bB 0 products

Trial Initial [A] (M) Initial [B] (M) Initial Rate (mol/L�s)

1 0.10 0.30 7.20 � 10�3

2 0.10 0.60 1.44 � 10�2

3 0.20 0.90 8.64 � 10�2

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Reaction RatesReaction Rates

Section 16.1 A Model for Reaction RatesIn your textbook, read about expressing reaction rates and explaining reactions andtheir rates.

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

According to the (1) , atoms, ions, and molecules must collide in

order to react. Once formed, the (2) is a temporary, unstable

arrangement of atoms that may then form products or may break apart to reform the reactants.

Every chemical reaction requires energy, and the minimum amount of energy that reacting

particles must have to form the activated complex is the (3) . In a

chemical reaction, the (4) is the change in concentration of a reactant

or product per unit time. It may be expressed using the units of (5) .

Use the energy diagram for the rearrangement reaction of methyl isonitrile to acetoni-trile to answer the following questions.

6. What kind of reaction is represented by this diagram,endothermic or exothermic?

7. What is the chemical structure identified at the top ofthe curve on the diagram?

8. What does the symbol Ea represent?

9. What does the symbol �E represent?


collision theory activated complex mol/(L�s)

activation energy reaction rate

Reaction Progress

H3C�N�C

H3C�C�N

a

Ener

gy

�E

Ea

H3C...C

N

�

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

Column A Column B

10. Expresses the average rate of loss of a reactant a. average reaction rate

11. Expressed as �quantity/�time b. positive number

12. Expresses the average rate of formation of a product c. negative number

Use the figure below to answer the following questions.

13. What molecules collided in collisions A, B, and C?

14. What do the arrows represent?

15. Which collision(s) formed products? What were the products?

16. Explain why the other collision(s) did not form products.

17. Which collision(s) formed an activated complex? Identify the activated complex.

Incorrect orientationRebound

Correct orientationCollisionCollision Activated complex Products

Correct orientationInsufficient energy

ReboundCollision

�A B

C

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Section 16.2 Factors Affecting Reaction RatesIn your textbook, read about the factors that affect reaction rates (reactivity, concentra-tion, surface, area, temperature, and catalysts).

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

1. Decreasing the concentration of reactants increases the collisionfrequency between reacting particles.

2. A heterogeneous catalyst exists in a different physical state thanthe reaction it catalyzes.

3. Increasing the concentration of a substance increases the kineticenergy of the particles that make up the substance.

4. Catalysts increase the rates of chemical reactions by raising theactivation energy of the reactions.

5. Increasing the surface area of a reactant increases the rate of thereaction.

6. Raising the temperature of a reaction increases the rate of thereaction by increasing the energy of the collisions betweenreacting particles.


7. A chemist heated a sample of steel wool in a burner flame exposed to oxygen in the air.He also heated a sample of steel wool in a container of nearly 100% oxygen. The steel-wool sample in the container reacted faster than the other sample. Explain why.

8. Would the chemist have observed the same results if he used a block of steel instead ofsteel wool? Explain your answer.

9. How would the reaction have differed if the steel wool was not heated?


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Section 16.3 Reaction Rate LawsIn your textbook, read about reaction rate laws and determining reaction order.

Use each of the terms below to complete the statements.

Equation 1 aA + bB 0 cC + dD

Equation 2 � � k[A]m[B]n

1. Equation 1 describes a .

2. Equation 2 expresses the mathematical relationship between the rate of a chemical

reaction and the concentrations of the reactants. This is known as the

.

3. The variable k in equation 2 is the , a numerical value that relates

the reaction rate and the concentration at a given temperature.

4. The variables m and n are the . These define how the rate is affected

by the concentrations of the reactants.

5. The square brackets [ ] represent .

6. The variable t represents .

Answer the questions about the following rate law.

Rate � k [A]1[B]2

7. What is the reaction order with respect to A?

8. What is the reaction order with respect to B?

9. What is the overall reaction order for the rate law?

10. Doubling the concentration of A will cause the rate to double. What would happen if you dou-

bled the concentration of B?

11. A reaction rate can be expressed as Rate 5 k[A]2. What is the reaction order for this reaction?

�[A]�

�t

chemical reaction rate law specific rate constant

reaction orders concentration time


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Section 16.4 Instantaneous Reaction Rates and ReactionMechanismsIn your textbook, read about instantaneous reaction rates.

Circle the letter of the choice that best completes the statement.

1. is determined by finding the slope of the straight line tangent to the curve of aplot of the change in concentration of a reactant versus time.

a. Instantaneous rate c. Reaction mechanism

b. Change in temperature d. Reaction order

2. A(n) consists of two or more elementary steps.

a. complex reaction c. reaction mechanism

b. elementary step d. reaction order

3. A(n) is a substance produced in an elementary step and consumed in anotherelementary step.

a. instantaneous rate c. reaction mechanism

b. intermediate d. rate-determining step

4. A(n) is the complete sequence of elementary reactions that make up a complexreaction.

a. instantaneous rate c. reaction mechanism

b. elementary step d. reaction order

5. The is the slowest of the elementary steps in a complex reaction.

a. instantaneous rate c. rate-determining step

b. intermediate d. reaction order

6. The can be used to determine the instantaneous rate for a chemical reaction.

a. rate-determining step c. products

b. intermediates d. rate law

7. An element or compound that reacts in one step of a complex reaction and reforms in aanother step of the complex reaction is

a. an intermediate.

b. a catalyst.

c. not part of the reaction mechanism.

d. shown in the net chemical equation for the reaction.


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8. To determine the instantaneous rate, you must know the specificrate constant, the concentrations of the reactants, and the reactionorders for the reaction.

9. A reaction rate that is defined as k[A][B] and that has a specificrate constant of 1.0 � 101 L/(mol�s), [A] � 0.1M, and [B] � 0.1M would have an instantaneous rate of 0.01 mol/(L�s).

In your textbook, read about reaction mechanisms.

Answer the following questions about the proposed reaction mechanism for the complexreaction below.

2NO(g) + 2H2(g) 0 N2(g) + 2H2O(g)

Proposed Mechanism 2NO 0 N2O2 (fast)

N2O2 + H2 0 N2O + H2O (slow)

N2O + H2 0 N2 + H2O (fast)

10. How many elementary steps make up the complex reaction?

11. What is the rate-determining step for this reaction?

12. What are N2O2 and N2O in the reaction?

13. Is there a catalyst involved in the reaction? Explain your answer.

14. What can you conclude about the activation energy for the rate-determining step?

15. If you wanted to increase the rate of the overall reaction, what would you do?



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CHAPTER 16

Assessment



1. 4. 7. 10.

2. 5. 8.

3. 6. 9.

Short Answer


11.

12.

Extended Response


13.

14.

15.


16. 18.

17. 19.


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