Chemistry: Matter and Change Challenge Problems

Challenge Problems

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CHALLENGE PROBLEMS

Contents To the Teacher . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . iv Chapter 1

Production of Chlorofluorocarbons, 1950–1992 . . . . . . . . . 1

Chapter 2

Population Trends in the United States . . . . . . . . . . . . . . . . 2

Chapter 3

Physical and Chemical Changes . . . . . . . . . . . . . . . . . . . . . 3

Chapter 4

Isotopes of an Element . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4

Chapter 5

Quantum Numbers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5

Chapter 6

Döbereiner’s Triads . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 6

Chapter 7

Abundance of the Elements . . . . . . . . . . . . . . . . . . . . . . . . 7

Chapter 8

Comparing the Structures of Atoms and Ions . . . . . . . . . . . 8

Chapter 9

Exceptions to the Octet Rule . . . . . . . . . . . . . . . . . . . . . . . . 9

Chapter 10 Balancing Chemical Equations . . . . . . . . . . . . . . . . . . . . . 10 Chapter 11 Using Mole-Based Conversions . . . . . . . . . . . . . . . . . . . . 11 Chapter 12 Mole Relationships in Chemical Reactions . . . . . . . . . . . . 12 Chapter 13 Intermolecular Forces and Boiling Points . . . . . . . . . . . . . 13 Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

Chapter 14 A Simple Mercury Barometer . . . . . . . . . . . . . . . . . . . . . . 14 Chapter 15 Vapor Pressure Lowering . . . . . . . . . . . . . . . . . . . . . . . . . 15 Chapter 16 Standard Heat of Formation . . . . . . . . . . . . . . . . . . . . . . . 16 Chapter 17 Determining Reaction Rates . . . . . . . . . . . . . . . . . . . . . . . 17 Chapter 18 Changing Equilibrium Concentrations in a Reaction . . . . . 18 Chapter 19 Swimming Pool Chemistry . . . . . . . . . . . . . . . . . . . . . . . . 19 Chapter 20 Balancing Oxidation–Reduction Equations . . . . . . . . . . . . 20 Chapter 21 Effect of Concentration on Cell Potential . . . . . . . . . . . . . 21 Chapter 22 Structural Isomers of Hexane . . . . . . . . . . . . . . . . . . . . . . 22 Chapter 23 Boiling Points of Organic Families . . . . . . . . . . . . . . . . . . 23 Chapter 24 The Chemistry of Life . . . . . . . . . . . . . . . . . . . . . . . . . . . 24 Chapter 25 The Production of Plutonium-239 . . . . . . . . . . . . . . . . . . . 25 Chapter 26 The Phosphorus Cycle . . . . . . . . . . . . . . . . . . . . . . . . . . . 26 Answer Key . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . T27

Challenge Problems

Chemistry: Matter and Change

iii

Name

CHAPTER

Date

1

Class

CHALLENGE PROBLEMS

C

hlorofluorocarbons (CFCs) were first produced in the laboratory in the late 1920s. They did not become an important commercial product until some time later. Eventually, CFCs grew in popularity until their effect on the ozone layer was discovered in the 1970s. The graph shows the combined amounts of two important CFCs produced between 1950 and 1992. Answer the following questions about the graph.

Amount of CFCs (billion kilograms)

Production of Chlorofluorocarbons, 1950–1992 400 350 300 250 200 150 100 50 0 1950

1960

Use with Chapter 1, Section 1.1

1970 Year

1980

1990

1. What was the approximate amount of CFCs produced in 1950? In 1960? In 1970?

2. In what year was the largest amount of CFCs produced? About how much was produced

that year?

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

3. During what two-year period did the production of CFCs decrease by the greatest

amount? By about how much did their production decrease?

4. During what two-year period did the production of CFCs increase by the greatest

amount? What was the approximate percent increase during this period?

5. How confident would you feel about predicting the production levels of CFCs during the

odd numbered years 1961, 1971, and 1981? Explain.

6. Could the data in the graph be presented in the form of a circle graph? Explain.

Challenge Problems

Chemistry: Matter and Change • Chapter 1

1

Name

Date

CHAPTER

2

Class

CHALLENGE PROBLEMS

Population Trends in the United States

Use with Chapter 2, Section 2.4

T

he population of the United States is becoming more diverse. The circle graphs below show the distribution of the U.S. population among five ethnic groups in 1990 and 2000. The estimated total U.S. population for those two years was 2.488  108 in 1990 and 2.754  108 in 2000. U.S. Population Distribution African American 11.8% Hispanic American 9.0% Asian American 2.8% Native American 0.70%

1990

2000 African American 12.2% Hispanic American 11.8%

Caucasian 75.7%

Asian American 3.8% Native American 0.70%

Caucasian 71.4%

(Percentages may not add up to 100% due to rounding.)

1. By how much did the total U.S. population increase between 1990 and 2000? What was

2. Calculate the total population for each of the five groups for 1990 and 2000.

3. Make a bar graph that compares the population for the five groups in 1990 and 2000. In

what ways is the bar graph better than the circle graphs? In what way is it less useful?

2

Chemistry: Matter and Change • Chapter 2

Challenge Problems

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

the percent increase during this period?

Name

CHAPTER

Date

3

Class

CHALLENGE PROBLEMS

Physical and Chemical Changes

Use with Chapter 3, Section 3.2

P

hysical and chemical changes occur all around us. One of the many places in which physical and chemical changes occur is the kitchen. For example, cooking spaghetti in a pot of water on the stove involves such changes. For each of the changes described below, tell (a) whether the change that occurs is physical or chemical, and (b) how you made your choice between these two possibilities. If you are unable to decide whether the change is physical or chemical, tell what additional information you would need in order to make a decision. 1. As the water in the pot is heated, its temperature rises.

2. As more heat is added, the water begins to boil and steam is produced.

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

3. The heat used to cook is produced by burning natural gas in the stove burner.

4. The metal burner on which the pot rests while being heated becomes red as its

temperature rises.

5. After the flame has been turned off, a small area on the burner has changed in color from

black to gray.

6. A strand of spaghetti has fallen onto the burner, where it turns black and begins to

smoke.

7. When the spaghetti is cooked in the boiling water, it becomes soft.

Challenge Problems

Chemistry: Matter and Change • Chapter 3

3

Name

Date

CHAPTER

4

Class

CHALLENGE PROBLEMS

Isotopes of an Element mass spectrometer is a device for separating atoms and molecules according to their mass. A substance is first heated in a vacuum and then ionized. The ions produced are accelerated through a magnetic field that separates ions of different masses. The graph below was produced when a certain element (element X) was analyzed in a mass spectrometer. Use the graph to answer the questions below.

30 Percent abundance

A

Use with Chapter 4, Section 4.3

25 20 15 10 5 0 190 192 194 196 198 200 202 204 206 208 210

Atomic mass (amu)

1. How many isotopes of element X exist? 2. What is the mass of the most abundant isotope? 3. What is the mass of the least abundant isotope? 4. What is the mass of the heaviest isotope? 5. What is the mass of the lightest isotope?

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

6. Estimate the percent abundance of each isotope shown on the graph.

7. Without performing any calculations, predict the approximate atomic mass for element

X. Explain the basis for your prediction.

8. Using the data given by the graph, calculate the weighted average atomic mass of

element X. Identify the unknown element.

4

Chemistry: Matter and Change • Chapter 4

Challenge Problems

Name

Date

5

CHAPTER

Class

CHALLENGE PROBLEMS

Quantum Numbers

Use with Chapter 5, Section 5.2

T

he state of an electron in an atom can be completely described by four quantum numbers, designated as n, , m, and ms. The first, or principal, quantum number, n, indicates the electron’s approximate distance from the nucleus. The second quantum number, , describes the shape of the electron’s orbit around the nucleus. The third quantum number, m, describes the orientation of the electron’s orbit compared to the plane of the atom. The fourth quantum number, ms, tells the direction of the electron’s spin (clockwise or counterclockwise).

The Schrödinger wave equation imposes certain mathematical restrictions on the quantum numbers. They are as follows: n can be any integer (whole number),

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

 can be any integer from 0 to n  1, m can be any integer from  to , and 1 or  1 ms can be    2 2 As an example, consider electrons in the first energy level of an atom, that is, n  1. In this case,  can have any integral value from 0 to (n  1), or 0 to (1  1). In other words,  must be 0 for these electrons. Also, the only value that m can have is 0. The electrons in 1 or  1 for m . These restrictions agree with the this energy level can have values of    s 2 2 observation that the first energy level can have only two electrons. Their quantum numbers 1 and 1, 0, 0  1 . are 1, 0, 0,    2 2 Use the rules given above to complete the table listing the quantum numbers for each electron in a boron atom. The correct quantum numbers for one electron in the atom is provided as an example. Boron (B) Electron

n



m

ms

1

1

0

0

1  2

2 3 4 5

Challenge Problems

Chemistry: Matter and Change • Chapter 5

5

Name

Date

CHAPTER

6

Class

CHALLENGE PROBLEMS

Döbereiner’s Triads

Use with Chapter 6, Section 6.2

O

ne of the first somewhat successful attempts to arrange the elements in a systematic way was made by the German chemist Johann Wolfgang Döbereiner (1780–1849). In 1816, Döbereiner noticed that the then accepted atomic mass of strontium (50) was midway between the atomic masses of calcium (27.5) and barium (72.5). Note that the accepted atomic masses for these elements today are very different from their accepted atomic masses at the time Döbereiner made his observations. Döbereiner also observed that strontium, calcium, and barium showed a gradual gradation in their properties, with the values of some of strontium’s properties being about midway between the values of calcium and barium. Döbereiner eventually found four other sets of three elements, which he called triads, that followed the same pattern. In each triad, the atomic mass of the middle element was about midway between the atomic masses of the other two elements. Unfortunately, because Döbereiner’s system did not turn out to be very useful, it was largely ignored.

Set 1 Element

Melting Point (°C) 219.6

Fluorine Chlorine

Set 2

Calculated:

Element

Actual:

Boiling Point (°C)

Krypton

153 Calculated:

Element

Tin

Actual:

6

Calcium

Lead

Chemistry: Matter and Change • Chapter 6

Calculated:

Strontium

1384

Set 6

Melting Point (°C) 937 Calculated:

Element

Boiling Point (°C)

Beryllium Magnesium

Actual: 62

1107

Actual: 39.098

Germanium

Boiling Point (°C)

Magnesium

Set 5

Element

Radon

Calculated:

Potassium

Set 4

Xenon

6.941

Element

Actual: 7.2

Bromine

Atomic Mass

Lithium Sodium

Set 3

1285 Calculated: Actual:

327

Calcium

851

Challenge Problems

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

Had Döbereiner actually discovered a way of identifying trends among the elements? Listed below are six three-element groups in which the elements in each group are consecutive members of the same group in the periodic table. The elements in each set show a gradation in their properties. Values for the first and third element in each set are given. Determine the missing value in each set by calculating the average of the two given values. Then, compare the values you obtained with those given in the Handbook of Chemistry and Physics. Record the actual values below your calculated values. Is the value of the property of the middle element in each set midway between the values of the other two elements in the set?

Name

CHAPTER

Date

7

Class

CHALLENGE PROBLEMS

Abundance of the Elements

Use with Chapter 7, Section 7.1

T

he abundance of the elements differs significantly in various parts of the universe. The table below lists the abundance of some elements in various parts of the universe. Use the table to answer the following questions. Abundance (Number of atoms per 1000 atoms)* Element

Universe

Hydrogen

927

Helium

71.8

Solar System

Earth

863

Earth’s Crust

Human Body

30

606

610

257

135

Oxygen

0.510

0.783

500

Nitrogen

0.153

0.0809

24

Carbon

0.0811

0.459

106

Silicon

0.0231

0.0269

140

210

Iron

0.0139

0.00320

170

19

* An element is not abundant in a region that is left blank.

1. What percent of all atoms in the universe are either hydrogen or helium? What percent of

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

all atoms in the solar system are either hydrogen or helium?

2. Explain the relatively high abundance of hydrogen and helium in the universe compared

to their relatively low abundance on Earth.

3. Only the top four most abundant elements on Earth and in Earth’s crust are shown in the

table. Name two additional elements you would expect to find among the top ten elements both on Earth and in Earth’s crust. Explain your choices.

4. Name at least three elements in addition to those shown in the table that you would

expect to find in the list of the top ten elements in the human body. Explain your choices.

Challenge Problems

Chemistry: Matter and Change • Chapter 7

7

Name

Date

CHAPTER

8

Class

CHALLENGE PROBLEMS

Comparing the Structures of Atoms and Ions

Use with Chapter 8, Section 8.1

T

he chemical properties of an element depend primarily on its number of valence electrons in its atoms. The noble gas elements, for example, all have similar chemical properties because the outermost energy levels of their atoms are completely filled. The chemical properties of ions also depend on the number of valence electrons. Any ion with a complete outermost energy level will have chemical properties similar to those of the noble gas elements. The fluoride ion (F), for example, has a total of ten electrons, eight of which fill its outermost energy level. F has chemical properties, therefore, similar to those of the noble gas neon. Shown below are the Lewis electron dot structures for five elements: sulfur (S), chlorine (Cl), argon (Ar), potassium (K), and calcium (Ca). Answer the questions below about these structures. S

Cl

Ar

K

Ca

1. Write the atomic number for each of the five elements shown above.

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

2. Write the electron configuration for each of the five elements.

3. Which of the above Lewis electron dot structures is the same as the Lewis electron dot

structure for the ion S2? Explain your answer.

4. Which of the above Lewis electron dot structures is the same as that for the ion Cl?

Explain your answer.

5. Which of the above Lewis electron dot structures is like that for the ion K? Explain

your answer.

6. Name an ion of calcium that has chemical properties similar to those of argon. Explain

your answer.

8

Chemistry: Matter and Change • Chapter 8

Challenge Problems

Name

CHAPTER

Date

9

Class

CHALLENGE PROBLEMS

Exceptions to the Octet Rule

Use with Chapter 9, Section 9.3

T

he octet rule is an important guide to understanding how most compounds are formed. However, there are a number of cases in which the octet rule does not apply. Answer the following questions about exceptions to the octet rule. 1. Draw the Lewis structure for the compound BeF2.

2. Does BeF2 obey the octet rule? Explain.

3. Draw the Lewis structure for the compound NO2.

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

4. Does NO2 obey the octet rule? Explain.

5. Draw the Lewis structure for the compound N2F2.

6. Does N2F2 obey the octet rule? Explain.

7. Draw the Lewis structure for the compound IF5.

8. Does IF5 obey the octet rule? Explain.

Challenge Problems

Chemistry: Matter and Change • Chapter 9

9

Name

Date

CHAPTER

10

Class

CHALLENGE PROBLEMS

Balancing Chemical Equations

Use with Chapter 10, Section 10.1

E

ach chemical equation below contains at least one error. Identify the error or errors and then write the correct chemical equation for the reaction.

1. K(s)  2H2O(l) 0 2KOH(aq)  H2(g)

2. MgCl2(aq)  H2SO4(aq) 0 Mg(SO4)2(aq)  2HCl(aq)

3. AgNO3(aq)  H2S(aq) 0 Ag2S(aq)  HNO3(aq)

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

4. Sr(s)  F2(g) 0 Sr2F

5. 2NaHCO3(s)  2HCl(aq) 0 2NaCl(s)  2CO2(g)

6. 2LiOH(aq)  2HBr(aq) 0 2LiBr(aq)  2H2O

7. NH4OH(aq)  KOH(aq) 0 KOH(aq)  NH4OH(aq)

8. 2Ca(s)  Cl2(g) 0 2CaCl(aq)

9. H2SO4(aq)  2Al(NO3)3(aq) 0 Al2(SO4)3(aq)  2HNO3(aq)

10

Chemistry: Matter and Change • Chapter 10

Challenge Problems

Name

Date

11

CHAPTER

Class

CHALLENGE PROBLEMS

Using Mole-Based Conversions

Use with Chapter 11, Section 11.3

T

he diagram shows three containers, each of which holds a certain mass of the substance indicated. Complete the table below for each of the three substances.

UF6 (g)

CCl3CF3 (l)

Pb (s)

225.0 g

200.0 g

250.0 g

Substance

Mass (g)

Molar Mass (g/mol)

Number of Moles (mol)

Number of Representative Particles

UF6(g) CCl3CF3(l)

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

Pb(s)

1. Compare and contrast the number of representative particles and the mass of UF6 with

the number of representative particles and mass of CCl3CF3. Explain any differences you observe.

2. UF6 is a gas used in the production of fuel for nuclear power plants. How many moles of

the gas are in 100.0 g of UF6?

3. CCl3CF3 is a chlorofluorocarbon responsible for the destruction of the ozone layer in

Earth’s atmosphere. How many molecules of the liquid are in 1.0 g of CCl3CF3?

4. Lead (Pb) is used to make a number of different alloys. What is the mass of lead present

in an alloy containing 0.15 mol of lead?

Challenge Problems

Chemistry: Matter and Change • Chapter 11

11

Name

Date

CHAPTER

12

Class

CHALLENGE PROBLEMS

Mole Relationships in Chemical Reactions

Use with Chapter 12, Section 12.2

T

he mole provides a convenient way of finding the amounts of the substances in a chemical reaction. The diagram below shows how this concept can be applied to the reaction between carbon monoxide (CO) and oxygen (O2), shown in the following balanced equation. 2CO(g)  O2(g) 0 2CO2(g) Use the equation and the diagram to answer the following questions. Moles of CO

3 Particles of CO

1

6

2 4 Grams of CO

Moles of CO2

5

7 Particles of CO2

Grams of CO2

1. What information is needed to make the types of conversions shown by double-arrow 1

in the diagram?

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

2. What conversion factors would be needed to make the conversions represented by

double-arrow 2 in the diagram for CO? By double-arrow 6 for CO2?

3. What information is needed to make the types of conversions represented by

double-arrows 3 and 7 in the diagram?

4. What conversion factors would be needed to make the conversions represented by

double-arrow 3 in the diagram for CO?

5. Why is it not possible to convert between the mass of a substance and the number of

representative particles, as represented by double-arrow 4 of the diagram?

6. Why is it not possible to use the mass of one substance in a chemical reaction to find the mass

of a second substance in the reaction, as represented by double-arrow 5 in the diagram?

12

Chemistry: Matter and Change • Chapter 12

Challenge Problems

Name

CHAPTER

Date

13

Class

CHALLENGE PROBLEMS

Intermolecular Forces and Boiling Points he boiling points of liquids depend partly on the mass of the particles of which they are made. The greater the mass of the particles, the more energy is needed to convert a liquid to a gas, and, thus, the higher the boiling point of the liquid. This pattern may not hold true, however, when there are significant forces between the particles of a liquid. The graph plots boiling point versus molecular mass for group 4A and group 6A hydrides. A hydride is a binary compound containing hydrogen and one other element. Use the graph to answer the following questions.

100 Boiling point (°C)

T

Use with Chapter 13, Section 13.3

H2O

H2Te

0

H2Se H2S

100

0 0

Group 6A hydrides

SiH4 CH4

SnH4 GeH4 Group 4A hydrides

50 100 Molecular mass

150

1. How do the boiling points of the group 4A hydrides change as the molecular masses of

the hydrides change?

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

2. What are the molecular structure and polarity of the four group 4A hydrides?

3. Predict the strength of the forces between group 4A hydride molecules. Explain how

those forces affect the boiling points of group 4A hydrides.

4. How do the boiling points of the group 6A hydrides change as the molecular masses of

the hydrides change?

5. What are the molecular structure and polarity of the four group 6A hydrides?

6. Use Table 9-4 in your textbook to determine the difference in electronegativities of the

bonds in the four group 6A hydrides.

Challenge Problems

Chemistry: Matter and Change • Chapter 13

13

Name

Date

CHAPTER

14

Class

CHALLENGE PROBLEMS

A Simple Mercury Barometer

I

n Figure 1, a simple mercury barometer is made by filling a long glass tube with mercury and then inverting the open end of the tube into a bowl of mercury. Answer the following questions about the simple mercury barometer shown here.

Use with Chapter 14, Section 14.1

Glass tube Mercury column

1. What occupies the space above the mercury column in the

Bowl of mercury

barometer’s glass tube?

At sea level

At 500 meters above sea level

Figure 1

Figure 2

2. What prevents mercury from flowing out of the glass tube into the bowl of mercury?

3. When the barometer in Figure 1 is moved to a higher elevation, such as an altitude of

4. Suppose the barometer in Figure 1 was carried into an open mine 500 meters below sea

level. How would the height of the mercury column change? Explain why.

5. Suppose the liquid used to make the barometer was water instead of mercury. How would

this substitution affect the barometer? Explain.

6. Suppose a tiny crack formed at the top of the barometer’s glass tube. How would this

event affect the column of mercury? Explain why.

14

Chemistry: Matter and Change • Chapter 14

Challenge Problems

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

5000 meters, the column of mercury changes as shown in Figure 2. Why is the mercury column lower in Figure 2 than in Figure 1?

Name

CHAPTER

Date

15

Class

CHALLENGE PROBLEMS

Vapor Pressure Lowering

Use with Chapter 15, Section 15.3

Y

ou have learned that adding a nonvolatile solute to a solvent lowers the vapor pressure of that solvent. The amount by which the vapor pressure is lowered can be calculated by means of a relationship discovered by the French chemist François Marie Raoult (1830–1901) in 1886. According to Raoult’s law, the vapor pressure of a solvent (P) is equal to the product of its vapor pressure when pure (P0) and its mole fraction (X) in the solution, or P  P0X The solution shown at the right was made by adding 75.0 g of sucrose (C12H22O11) to 500.0 g of water at a temperature of 20°C. Answer the following questions about this solution.

Solution

Water molecule

Sucrose molecule

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

1. Why do the sugar molecules in the solution lower the vapor pressure of the water?

2. What is the number of moles of sucrose in the solution?

3. What is the number of moles of water in the solution?

4. What is the mole fraction of water in the solution?

5. What is the vapor pressure of the solution if the vapor pressure of pure water at 20°C is

17.54 mm Hg?

6. How much is the vapor pressure of the solution reduced from that of water by the

addition of the sucrose?

Challenge Problems

Chemistry: Matter and Change • Chapter 15

15

Name

Date

CHAPTER

16

Class

CHALLENGE PROBLEMS

Standard Heat of Formation C(s)  O2(g)

H

H  110 kJ/mol CO(g) 

1 O (g) 2 2

H  393 kJ/mol Enthalpy

ess’s law allows you to determine the standard heat of formation of a compound when you know the heats of reactions that lead to the production of that compound. The first diagram on the right shows how Hess’s law can be used to calculate the heat of formation of CO2 by knowing the heats of reaction of two steps leading to the production of CO2. Use this diagram to help you answer the questions below about the second diagram.

Use with Chapter 16, Section 16.4

H  283 kJ/mol

The equations below show how NO2 can be formed in two ways: directly from the elements or in two steps. H  33 kJ/mol

1 1  N2(g)   O2(g) 0 NO(g) 2 2

H  91 kJ/mol

1 O (g) 0 NO (g) NO(g)   2 2 2

H  58 kJ/mol

CO2(g)

C

NO(g)  1/2 O2(g)

1. On the diagram at the right, draw arrowheads

to show the directions in which the three lines labeled 1, 2, and 3 should point. 2. Write the correct reactants and/or products on

2 H  58 kJ/mol

each of the lines labeled A, B, and C.

1 H  91 kJ/mol

each number on the diagram.

Enthalpy

3. Write the correct enthalpy change next to

B

NO2(g)

3 H  33 kJ/mol A

16

Chemistry: Matter and Change • Chapter 16

1/2 N2(g)  O2(g)

Challenge Problems

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

1  N2(g)  O2(g) 0 NO2(g) 2 or

Name

Date

CHAPTER

17

Class

CHALLENGE PROBLEMS

Determining Reaction Rates initrogen pentoxide decomposes to produce nitrogen dioxide and oxygen as represented by the following equation. 2N2O5(g) 0 4NO2(g)  O2(g) The graph on the right represents the concentration of N2O5 remaining as the reaction proceeds over time. Answer the following questions about the reaction.

1.6 Concentration (mol/L)

D

Use with Chapter 17, Section 17.1

1.4 1.2 1.0 0.8 0.6 0.4 0.2 0 0 1 2 3 4 5 6 7 8 9 10 Time (h)

1. What is the concentration of N2O5 at the beginning of the experiment? After 1 hour?

After 2 hours? After 10 hours?

2. By how much does the concentration of N2O5 change during the first hour of the

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

reaction? Calculate the percentage of change the concentration undergoes during the first hour of the reaction.

3. The instantaneous rate of reaction is defined as the change in concentration of reactant

during some specified time period, or instantaneous rate of reaction = [N2O5]/t. What is the instantaneous rate of reaction for the decomposition of N2O5 for the time period between the first and second hours of the reaction? Between the second and third hours? Between the sixth and seventh hours?

4. What is the instantaneous rate of reaction for the decomposition of N2O5 between the sec-

ond and fourth hours of the reaction? Between the third and eighth hours of the reaction?

5. How long does it take for 0.10 mol of N2O5 to decompose during the tenth hour of the reaction?

6. What is the average rate of reaction for the decomposition of N2O5 overall?

Challenge Problems

Chemistry: Matter and Change • Chapter 17

17

Name

Date

CHAPTER

18

Class

CHALLENGE PROBLEMS

R

eversible reactions eventually reach an equilibrium condition in which the concentrations of all reactants and products are constant. Equilibrium can be disturbed, however, by the addition or removal of either a reactant or product. The graph on the right shows how the concentrations of the reactants and product of a reaction change when equilibrium is disturbed. Use the graph to answer the following questions.

Concentration (mol/L)

Changing Equilibrium Concentrations in a Reaction 8 7 6 5 4 3 2 1 0

Use with Chapter 18, Section 18.1

SO2

SO2 O2

O2

SO3

SO3

0 1 2 3 4 5 6 7 8 9 10 Time (sec)

1. Write the equation for the reaction depicted in the graph.

2. Write the equilibrium constant expression for the reaction.

3. Explain the shapes of the curves for the three gases during the first 2 minutes of the

4. At approximately what time does the reaction reach equilibrium? How do you know

equilibrium has been reached?

5. What are the concentrations of the three gases at equilibrium?

6. Calculate the value of Keq for the reaction.

7. Describe the change made in the system 4 minutes into the reaction. Tell how you know

the change was made.

8. At what time does the system return to equilibrium?

18

Chemistry: Matter and Change • Chapter 18

Challenge Problems

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

reaction.

Name

CHAPTER

Date

19

Class

CHALLENGE PROBLEMS

Swimming Pool Chemistry

Use with Chapter 19, Section 19.2

T

he presence of disease-causing bacteria in swimming pools is a major health concern. Chlorine gas is added to the water in some large commercial swimming pools to kill bacteria. However, in most home swimming pools, either solid calcium hypochlorite (Ca(OCl)2) or an aqueous solution of sodium hypochlorite (NaOCl) is used to treat the water. Both compounds dissociate in water to form the weak acid hypochlorous acid (HOCl). Hypochlorous acid is a highly effective bactericide. By contrast, the hypochlorite ion (OCl) is not a very effective bactericide. Use the information above to answer the following questions about the acid-base reactions that take place in swimming pools. 1. Write an equation that shows the reaction between hypochlorous acid and water. Identify

the acid, base, conjugate acid, and conjugate base in this reaction.

2. Write an equation that shows the reaction that occurs when the hypochlorite ion (OCl),

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

in the form of calcium hypochlorite or sodium hypochlorite, is added to water. Name the acid, base, conjugate acid, and conjugate base in this reaction.

3. What effect does the addition of hypochlorite ion have on the pH of swimming pool water?

4. The effectiveness of hypochlorite ion as a bactericide depends on pH. How does high pH

affect the equilibrium reaction described in question 2? What effect would high pH have on the bacteria?

5. In the presence of sunlight, hypochlorite ion decomposes to form chloride ion and oxygen

gas. Write an equation for this reaction and tell how it affects the safety of pool water.

Challenge Problems

Chemistry: Matter and Change • Chapter 19

19

Name

Date

CHAPTER

20

Class

CHALLENGE PROBLEMS

Balancing Oxidation– Reduction Equations

Use with Chapter 20, Section 20.3

S

cientists have developed a number of methods for protecting metals from oxidation. One such method involves the use of a sacrificial metal. A sacrificial metal is a metal that is more easily oxidized than the metal it is designed to protect. Galvanized iron, for example, consists of a piece of iron metal covered with a thin layer of zinc. When galvanized iron is exposed to oxygen, it is the zinc, rather than the iron, that is oxidized. Water heaters often contain a metal rod that is made by coating a heavy steel wire with magnesium or aluminum. In this case, the magnesium or aluminum is the sacrificial metal, protecting the iron casing of the heater from corrosion. The diagram shows a portion of a water heater containing a sacrificial rod. Answer the following questions about the diagram.

Steel wire Sacrificial metal

Iron casing

Water

1. In the absence of a sacrificial metal, oxygen dissolved in water may react with the iron

2. Balance the oxidation–reduction equation for this reaction:

Fe(s)  O2(aq)  H2O 0 Fe(OH)2(aq)

3. Write the two half-reactions for this example of corrosion.

4. Suppose the sacrificial rod in the diagram above is coated with aluminum metal. Write

the balanced equation for the reaction of aluminum with oxygen dissolved in the water. (Hint: The product formed is aluminum hydroxide (Al(OH)3).

5. Write the two half-reactions for this example of corrosion.

6. Suppose that some iron in the casing of the water heater is oxidized, as shown in the

equation of question 2 above. The sacrificial metal (aluminum, in this case) immediately restores the Fe2 ions to iron atoms. Write two half-reactions that represent this situation.

20

Chemistry: Matter and Change • Chapter 20

Challenge Problems

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

casing of the heater. One product formed is iron(II) hydroxide (Fe(OH)2). Which element is oxidized and which is reduced in this reaction?

Name

Date

21

CHAPTER

Class

CHALLENGE PROBLEMS

Effect of Concentration on Cell Potential

Use with Chapter 21, Section 21.1

I

n a voltaic cell where all ions have a concentration of 1M, the cell potential is equal to the standard potential. For cells in which ion concentrations are greater or less than 1M, as shown below, an adjustment must be made to calculate cell potential. That adjustment is expressed by the Nernst equation: [product ion]x 0.0592 log  Ecell  E 0cell    n [reactant ion]y

In this equation, n is the number of moles of electrons transferred in the reaction, and x and y are the coefficients of the product and reactant ions, respectively, in the balanced half-cell reactions for the cell.

Voltmeter

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

Ag

Ag 1.0  102M

Cu

Cu2 1.0  103M

1. Write the two half-reactions and the overall cell reaction for the cell shown above.

2. Use Table 21-1 in your textbook to determine the standard potential of this cell. 3. Write the Nernst equation for the cell.

4. Calculate the cell potential for the ion concentrations shown in the cell. Challenge Problems

Chemistry: Matter and Change • Chapter 21

21

Name

Date

CHAPTER

22

Class

CHALLENGE PROBLEMS

Structural Isomers of Hexane

Use with Chapter 22, Sections 22.1 and 22.3

T

he structural formula of an organic compound can sometimes be written in a variety of ways, but sometimes structural formulas that appear similar can represent different compounds. The structural formulas below are ten ways of representing compounds having the molecular formula C6H14. a. CH3 CH2

e. CH3 CH2

CH2

CH2

CH3

CH2 CH3 CH

i. CH2

CH2

CH3

CH3

CH2

CH2

CH3

CH3 CH3

CH2

CH

CH2

CH3

b. CH3 CH

f. CH3 CH2

CH2

CH3

CH

CH

CH3

CH3 CH3

j.

CH2

CH3

c.

g. CH2

CH3 CH3

CH

CH

CH3

CH3

CH CH2

CH3 CH3

CH3

h.

CH3 CH3

C

CH2

CH3

CH3 CH3

CH3

CH

CH2 CH2

CH3

1. In the spaces provided, write the correct name for each of the structural formulas, labeled

a–j, above. a.

e.

i.

b.

f.

j.

c.

g.

d.

h.

2. How many different compounds are represented by the structural formulas above? What

are their names?

22

Chemistry: Matter and Change • Chapter 22

Challenge Problems

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

d.

Name

CHAPTER

Date

23

Class

CHALLENGE PROBLEMS

Boiling Points of Organic Families he most important factor determining the boiling point of a substance is its atomic or molecular mass. In general, the larger the atomic or molecular mass of the substance, the more energy is needed to convert the substance from the liquid phase to the gaseous phase. As an example, the boiling point of ethane (molecular mass  30; boiling point  89°C) is much higher than the boiling point of methane (molecular mass  16; boiling point  161°C). Intermolecular forces between the particles of a liquid also can affect the liquid’s boiling point. The graph shows trends in the boiling points of four organic families: alkanes, alcohols, aldehydes, and ethers. Use the graph and your knowledge of intermolecular forces to answer the following questions.

100 Boiling point (°C)

T

Use with Chapter 23, Section 23.3

50 0 50

30

40

50 60 70 Molecular mass

 alkane  alcohol

80

 aldehyde  ether

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

1. For any one family, what is the relationship between molecular mass and boiling point?

2. For compounds of similar molecular mass, which family of the four shown in the graph

has the lowest boiling points? Which family has the highest boiling points?

3. Find and list the boiling points for ethanol (molecular mass  46) and dimethyl ether

(molecular mass  46) on the graph. Why would you expect these two compounds to have relatively similar boiling points?

4. Find the aldehyde with a molecular mass of about 58. Name that aldehyde and write its

chemical formula.

5. Can this aldehyde form hydrogen bonds? Can other aldehydes form hydrogen bonds?

Explain.

Challenge Problems

Chemistry: Matter and Change • Chapter 23

23

Name

Date

24

CHALLENGE PROBLEMS

The Chemistry of Life

Use with Chapter 24, Section 24.4

P

roteins are synthesized when RNA molecules translate the DNA language of nitrogen bases into the protein language of amino acids using a genetic code. The genetic code is found in RNA molecules called messenger RNA (mRNA), which are synthesized from DNA molecules. The genetic code consists of a sequence of three nitrogen bases in the mRNA, called a codon. Most codons code for specific amino acids. A few codons code for a stop in the synthesis of proteins. The table shows the mRNA codons that make up the genetic code. To use the table, read the three nitrogen bases in sequence. The first base is shown along the left side of the table. The second base is shown along the top of the table. The third base is shown along the right side of the table. For example, the sequence CAU codes for the amino acid histidine (His). The table gives abbreviations for the amino acids. Answer the following questions about the genetic code.

The Genetic Code Second base U

C

A

G

} }

C UCU Phe UCC UCA Leu UCG CCU CCC Leu CCA CCG ACU Ile ACC ACA Met ACG GCU GCC Val GCA GCG

A UAU UAC Ser UAA UAG CAU CAC Pro CAA CAG AAU AAC Thr AAA AAG GAU GAC Ala GAA GAG

G

}Tyr Stop Stop

} His } Gln } Asn } Lys } Asp } Glu

UGU UGC UGA UGG CGU CGC CGA CGG AGU AGC AGA AGG GGU GGC GGA GGG

} Cys

} }

U C Stop A Trp G U C Arg A G U Ser C A Arg G U C Gly A G

Third base

First base

U

UUU UUC UUA UUG CUU CUC CUA CUG AUU AUC AUA AUG GUU GUC GUA GUG

1. What amino acid is represented by each of the following codons? a. CUG

b. UCA

2. Write the sequence of amino acids for which the following mRNA sequence codes.

-C-A-U-C-A-C-C-G-G-U-C-U-U-U-U-C-U-U-

3. Errors sometimes occur when mRNA molecules are synthesized from DNA molecules.

Nitrogen bases may be omitted, an extra nitrogen base may be added, or a nitrogen base may be changed during synthesis. The two mRNA sequences shown below are examples of such errors. In each case, tell how the mRNA sequence shown differs from the correct mRNA sequence given in question 2. a. -C-A-U-C-A-C-C-G-G-U-U-C-U-U-U-U-C-U-U-

b. -C-A-U-U-A-C-C-G-G-U-C-U-U-U-U-C-U-U-

4. Write the amino acid sequence for each of the mRNA sequences shown in question 3. a. b. 24

Chemistry: Matter and Change • Chapter 24

Challenge Problems

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

CHAPTER

Class

Name

CHAPTER

Date

25

Class

CHALLENGE PROBLEMS

The Production of Plutonium-239

Use with Chapter 25, Section 25.4

45p 75n

W

hen nuclear fission was first discovered, only two isotopes, uranium-233 and uranium-235, were known of being capable of undergoing this nuclear change. Scientists later discovered a third isotope, plutonium-239, also could undergo nuclear fission. Plutonium-239 does not occur in nature but can be made synthetically in nuclear reactors and particle accelerators.

92p 143n

1n 0

92p 143n

1n 0

1n 0

A

92p 146n

Source of neutrons

The diagram shows the process by which plutonium-239 is made in nuclear reactors. Answer the questions about the diagram.

C

0

0

1n 0 0 –1 0 –1

1. Identify the isotope whose nucleus is labeled A in the

diagram.

B

D

F 48p 77n

E G

2. Name the type of nuclear reaction that occurs when a

neutron strikes nucleus A.

Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

3. Identify the isotope whose nucleus is labeled B. 4. Besides fragmented nuclei, what else is produced when a neutron strikes nucleus A? 5. Identify the isotope whose nucleus is labeled C. 6. Write the nuclear equation for the reaction that occurs when a neutron strikes nucleus C.

Identify the product D formed in the reaction.

7. Write the nuclear equation for the decay of nucleus D. Identify isotope E formed in the

reaction.

8. Write a balanced nuclear equation for the decay of nucleus E. Identify isotope F formed

in the reaction.

9. Name the type of nuclear reaction that occurs when a neutron strikes nucleus F.

10. Write the nuclear equation for the reaction that occurs when a neutron strikes nucleus F.

Identify isotope G formed in the reaction.

Challenge Problems

Chemistry: Matter and Change • Chapter 25

25

Name

Date

CHAPTER

26

Class

CHALLENGE PROBLEMS

The Phosphorus Cycle

Use with Chapter 26, Section 26.4

P

hosphorus is an important element both in organisms and in the lithosphere. In organisms, phosphorus occurs in DNA and RNA molecules, cell membranes, bones and teeth, and in the energy–storage compound adenosine triphosphate (ATP). In the lithosphere, phosphorus occurs primarily in the form of phosphates, as a major constituent of many rocks and minerals. Phosphate rock is mined to produce many commercial products, such as fertilizers and detergents. When these products are used, phosphates are returned to the lithosphere and hydrosphere. Thus, phosphorus—like carbon and nitrogen—cycles in the environment. Use the diagram of the phosphorus cycle to answer the questions below. Phosphate rocks

Phosphate rocks Copyright © Glencoe/McGraw-Hill, a division of the McGraw-Hill Companies, Inc.

Geological uplift

1. By what methods does phosphorus get into soil?

2. By what method do plants obtain the phosphorus they need?

3. By what method do animals obtain the phosphorus they need?

4. In what way is the phosphorus cycle different from the carbon and nitrogen cycles you

studied in the textbook?

5. The phosphorus cycle has both short-term and long-term parts. Use different colored

pencils to show each part on the diagram.

26

Chemistry: Matter and Change • Chapter 26

Challenge Problems