Kelter P., Mosher M., Scott A. Chemistry. The Practical Science

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198 Chapter 5 Energy

landscape. The alternative strategy of harnessing energy from nuclear processes

was eagerly adopted in many countries in the latter half of the twentieth century.

The source of this energy will be discussed fully in Chapter 21.

One of the most attractive options is to make increasing use of

renewable

sources of energy

, which can be rapidly replaced by natural processes. This is in

contrast to nonrenewable sources, particularly the fossil fuels: coal, oil, and

natural gas. Staff at the U.S. National Renewable Energy Laboratory in Colorado

have studied the feasibility of using renewable energy sources to meet all of

America’s energy needs. The United States is the largest single user of energy in

the world, accounting for approximately 50% of global consumption. A feasible

plan to meet all or most of U.S. energy needs with renewable sources of energy

would encourage countries around the world to make a similar commitment.

According to the Renewable Energy Laboratory, the main renewable energy

sources that could meet the needs of the United States are

■

Solar cells (also known as photovoltaic systems, or PVs), which convert the

energy of sunlight into electricity.

■

Solar thermal systems, which covert the energy of sunlight into heat.

■

Biomass conversion, which releases energy from the chemical conversion—

often the burning—of plants and trees that can be grown as quickly as they

are consumed.

The Hoover Dam harnesses water to gen-

erate electricity.

Windmills can convert wind into electricity.

Solar energy can be used to make electricity.

■

Hydroelectric systems, which use the power of falling water is used to turn

turbines that generate electricity.

■

Wind power, which uses turbines driven by the force of the wind to generate

electricity.

■

Geothermal systems, which involve drilling deep into the earth and exploiting

the ﬂow of heat from the interior of the earth out toward the surface.

Each of these systems could readily make a signiﬁcant contribution to power

generation in the United States. They already contribute 8%, as shown in Fig-

ure 5.25, but that could rise toward 50% or even 100% if social and political in-

centives were in place. One of the most interesting points made in the Renewable

Energy Laboratory’s report is summarized in Figure 5.26, which shows the total

area that would have to be devoted to solar cells sufﬁcient to generate all the en-

ergy required by the United States. It is a large area, equivalent to more than 10%

of the area of the state of Nevada, but that is less than 25% of the area of the

country that is currently paved over with roads and streets. So equipped, the

United States could garner enough energy from the sun to meet all its needs.

Solar energy arrives free of charge, although providing the technology to cap-

ture, distribute, and use it entails signiﬁcant costs. The energy is also very “clean”

in the sense that its use does not release anywhere near the amount of pollutants

released by the burning of fossil fuels. Some pollution would inevitably be asso-

ciated with the manufacturing processes and other activities involved in making

and maintaining the solar cells, but the overall effect of a switch to solar is likely

to be a huge reduction in pollution. One of the main challenges of the future

might be to develop the chemical systems needed to make full use of the free en-

ergy that ﬂoods down on the planet every day from the Sun.

When the possible contributions from biomass, hydroelectric, wind, and

geothermal power are added to the equation, it all amounts to a persuasive argu-

ment for continuing to explore the use of “renewables”as an energy source for the

future.

5.6 Energy Choices 199

FIGURE 5.26

Total area required for a solar cell power

plant to meet the total U.S. annual elec-

trical power demand is represented by

the square on this map.

Photovoltaic cells can be used to power almost anything.

200 Chapter 5 Energy

The Bottom Line

■

Chemical changes are accompanied by the gain or

release of energy. (Section 5.1)

■

Chemicals can store huge amounts of energy and

then release it as soon as a chemical reaction begins.

(Section 5.1)

■

Thermodynamics is the study of energy changes and

exchanges. (Section 5.1)

■

Energy comes in two basic forms, which we call ki-

netic energy (the energy of motion) and potential

energy (positional energy). (Section 5.1)

■

Energy is never created or destroyed; it is only trans-

ferred from place to place and converted from one

form into another. This is the law of conservation of

energy. (Section 5.1)

■

A chemical reaction that releases energy from the

chemicals involved is known as an exothermic reac-

tion, because energy is ﬂowing out of the system and

into the surroundings. A reaction that absorbs en-

ergy into the chemicals involved is known as an en-

dothermic reaction. (Section 5.1)

■

All chemical reactions begin with an input of energy,

needed to “jolt” the chemicals into reacting. The en-

ergy required to make this happen is known as the

activation energy of the reaction. (Section 5.1)

■

The total change in the energy of a chemical system,

as it undergoes a chemical reaction, is equal to the

heat ﬂow (q), known as the heat of reaction, and the

work done (w): U = q + w. (Section 5.1)

■

The SI unit of energy is the joule, which we can re-

late to the more familiar units of calories and (in the

context of food) Calories. (Section 5.2)

■

Each substance has a particular speciﬁc heat capacity

(c), often simply called its speciﬁc heat, which is the

amount of heat needed to raise the temperature of

1 gram of the substance by 1 degree Celsius (or 1 kel-

vin) when the pressure is constant. (Section 5.3)

■

Hess’s law states that the enthalpy change of a chem-

ical reaction is independent of the chemical path or

mechanism involved in the reaction. (Section 5.5)

■

The standard enthalpy change for a reaction can be

calculated by subtracting the sum of the enthalpies

of formation of the reactants of the reaction from

the sum of the enthalpies of formation of the prod-

ucts of the reaction. (Section 5.5)

■

Making appropriate choices about our energy

sources will be an important part of building a suc-

cessful and sustainable future for humanity. (Sec-

tion 5.6)

Key Words

activation energy The energy that is required to initiate

a chemical reaction. (p. 174)

biomass conversion The release of energy from the

chemical conversion—often simply the burning—

of plants and trees. (p. 198)

bomb calorimeter Apparatus in which a chemical reac-

tion occurs in a closed container, allowing the energy

released or absorbed to be measured. (p. 184)

Calorie (Note capital C.) Unit of energy equal to 1000

calories (1 kilocalorie) and to 4184 joules. (p. 179)

calorie (Note lowercase c.) Unit of energy equal to

4.184 joules. (p. 180)

calorimeter An apparatus in which quantities of heat

can be measured. (p. 183)

calorimetry The study of the transfer of heat in a pro-

cess. (p. 183)

chemical energy Energy that is stored in a substance as

a result of the motions and positions of its atomic

nuclei and their electrons. (p. 172)

constant-volume calorimetry A form of calorimetry in

which the reacting system is sealed within a chamber

of ﬁxed volume, and the only way the system can re-

lease or gain energy is by the exchange of heat with

the surroundings. (p. 184)

electromagnetic radiation Energy that propagates through

space. Examples include visible light, X-rays, and

radiowaves. (p. 170)

electroweak force The fundamental force responsible for

the attraction between objects carrying opposite

electric charges, for the repulsion between objects

carrying the same electric charge, and for some trans-

formations within subatomic particles. It is also re-

sponsible for the phenomena of magnetism and light.

(p. 169)

endothermic reaction A reaction that absorbs energy

from the surroundings. (p. 175)

enthalpy A thermodynamic quantity symbolized by H

and deﬁned as H = U + PV.(p. 186)

exothermic reaction A reaction that releases energy into

the surroundings. (p. 174)

ﬁrst law of thermodynamics The total change in the

closed system’s energy in a chemical process is equal

Key Words 201

to the heat ﬂow (q) into the system and the work

done (w) on the system:

U = q + w

Energy is neither created nor destroyed; it is only

transferred from place to place and converted from

one form into another: U

system

+ U

surroundings

= 0.

(p. 177)

geothermal systems Power-generating systems that in-

volve drilling deep into the earth and exploiting the

ﬂow of heat from the interior of the Earth out toward

the surface. (p. 199)

gravitational force The fundamental force that causes

all objects with mass to be attracted to one another.

(p. 169)

heat (q) The energy that is exchanged between a system

and its surroundings because of a difference in tem-

perature between the two. (p. 174)

heat capacity The amount of heat needed to raise the

temperature of any particular amount of a substance

by 1°C. (p. 182)

heat of reaction The energy as heat released or absorbed

during the course of a reaction. (p. 186)

Hess’s law Thermodynamic law stating that the en-

thalpy change of a chemical reaction is independent

of the chemical path or mechanism involved in the

reaction. This enables us to determine the enthalpy

changes of reactions that might be very difﬁcult to

actually perform. (p. 192)

hydroelectric systems Power-generating systems in which

the power of falling water is used to generate electric-

ity. (p. 199)

internal energy (U) The energy of a system deﬁned as

the sum of the kinetic and potential energies. Ab-

solute internal energy is difﬁcult to measure. (p. 177)

joule (J) The SI unit of energy. In terms of base units,

1 Joule = 1 kg·m

·s

–2

.(p. 179)

kilocalorie Unit of energy equal to 1000 calories and to

1 Calorie (note capital C). (p. 180)

kinetic energy The energy things possess as a result of

their motion. (p. 170)

law of conservation of energy Law stating that energy is

neither created nor destroyed but is only transferred

from place to place or converted from one form into

another. (p. 171)

molar heat capacity The heat capacity of one mole of a

substance. (p. 182)

photovoltaic systems (PVs) Fabricated systems that con-

vert the energy of sunlight into electricity. Also

known as solar cells. (p. 198)

potential energy The energy things possess as a result of

their positions, such as their position in a gravita-

tional or electromagnetic ﬁeld. (p. 170)

reference form The most stable form of the element at

standard conditions. (p. 187)

renewable sources of energy Energy sources that can be

rapidly replaced by natural processes. (p. 198)

solar cells Fabricated systems that convert the energy of

sunlight into electricity. Also known as photovoltaic

systems or PVs. (p. 198)

solar thermal systems Fabricated systems that convert

the energy of sunlight into heat. (p. 198)

speciﬁc heat The amount of heat needed to raise the

temperature of one gram of a substance by one de-

gree Celsius (or one kelvin) when the pressure is con-

stant. Also known as the substance’s speciﬁc heat

capacity. (p. 181)

speciﬁc heat capacity (c) The amount of heat needed to

raise the temperature of one gram of a substance by

one degree Celsius (or one kelvin) when the pressure

is constant. Also known as the substance’s speciﬁc

heat. (p. 181)

standard enthalpy of combustion (

H°) The enthalpy

change when one mole of a substance in its standard

state is completely burned in oxygen gas.Also known as

the substance’s standard heat of combustion. (p. 189)

standard enthalpy of formation (

H°) The enthalpy

change for the formation of one mole of a substance

in its standard state from its elements in their refer-

ence form. Also known as the standard heat of for-

mation. (p. 188)

standard enthalpy of reaction (

rxn

H°) Enthalpy change of

a reaction in which all of the reactants and products

are in their standard states. (p. 187)

standard state The state of a chemical under a set of

standard conditions, usually at 1 atmosphere of pres-

sure and a concentration of exactly 1 molar for any

substances in solution. Standard states are often re-

ported at 25°C. (p. 187)

state function A property of a system that depends only

on its present state, not on the path by which it

reached that state. (p. 191)

strong nuclear force The force that holds protons and

neutrons together in atomic nuclei. (p. 169)

surroundings Everything in contact with a chemical sys-

tem. Energy may ﬂow between the surroundings and

system. Strictly speaking, the surroundings of a chem-

ical system comprise the entire universe except the

system. (p. 168)

system Any set of chemicals whose energy change we

are interested in. (p. 168)

thermochemistry The study of energy exchange in chem-

ical processes. (p. 168)

thermodynamics The study of energy exchange. (p. 168)

universe The space consisting of both the system and

the surroundings. (p. 168)

wind power The use of the energy of the wind to gener-

ate electricity. (p. 199)

work (w) Force acting on an object over a distance.

(p. 169)

Gas in

cylinder

14.02 kJ

9.33 kJ

Piston

moves

Gas in

cylinder

67 kJ

23 kJ

Piston

moves

202 Chapter 5 Energy

Focus Your Learning

The answers to the odd-numbered problems and some selected

problems appear at the back of the book, as represented by the

blue numbering.

Section 5.1 The Concept of Energy

Skill Review

1. Imagine a skateboarder at the top of a half-pipe. Describe the

changes in potential energy and in kinetic energy as the skater

goes from the top of one side of the half-pipe to the top of the

other side.

2. Describe the changes in potential and kinetic energyof a snow-

boarder going over a hill. If there were another uphill slope

waiting at the bottom of the ﬁrst hill, would the snowboarder’s

potential energy increase or decrease as she or he came down

the ﬁrst hill and immediately began to climb the second?

3. Calculate the kinetic energy of the following:

a. A 185-lb chemistry professor jogging at 8.0 mi/h.

b. A 42-g hackey sack being kicked at 0.78 m/s.

c. A molecule of CO

moving at 560 m/s.

4. Calculate the kinetic energy of the following:

a. A 2-ton truck lumbering along at 60.0 mi/h.

b. A 200-g apple falling from a tree at 20 mi/h.

c. A 0.25-kg ball moving at 1.75 m/s.

5. Describe the following atomic phenomena as primarily a

function of the potential or kinetic energy of the system.

a. The attractive force between an electron in a hydrogen

atom and its nucleus.

b. The vibration of two hydrogen atoms bonded in a

molecule of H

c. The movement of a hydrogen molecule inside a balloon.

6. Describe the following atomic phenomena as primarily a

function of the potential or kinetic energy of the system.

a. The attraction of a hydrogen atom’s electron to oxygen in

water.

b. The repulsion of a hydrogen atom’s electron from the

other hydrogen electron in H

c. The collision of two molecules as they react to generate a

product.

7. Review the deﬁnition of heat. What is being transferred when

heat moves between a system and its surroundings? How is

this transfer accomplished?

8. What are the three main forms in which energy can be trans-

ferred between a system and its surroundings? Give an exam-

ple of each.

9. When gasoline in an automobile engine undergoes combus-

tion, new compounds and heat are produced. Identify the

system and the surroundings in this process. Is the system

gaining or losing energy in this process? What is the sign on

w for this process?

10. When a cold pack is applied to a sports injury, the athlete

may remark that the pack feels very cold. Identify the system

and surroundings associated with the cold pack. Is the system

gaining or losing energy in this process? What is the sign on q

for this process?

11. In a chemical reaction, the value of q was determined to

be +24 joules (J). The work function was determined to be

+12 J. What is the total energy change for this system? Did

the surroundings gain or lose energy?

12. In a chemical reaction, the value of q was determined to

be −200 joules (J). The work function was determined to be

+158 J. What is the total energy change for this system? Did

the surroundings gain or lose energy?

13. Calculate the total energychange for each of thefollowing sys-

tems.Do the surroundings gain or lose energyin each process?

a. q =+45 J; w =+45 J

b. q =−266 J; w = 1.2 kJ

c. q = 23.4 kJ; w =−14 kJ

14. Calculate the total energy change for each of the following

systems. Does the system gain or lose energy in each process?

a. q =−23 J; w =−37 J

b. q =−88 J; w =+36 J

c. q =+105 J; w =−133 J

Focus Your Learning 203

Chemical Applications and Practices

15. Burning hydrocarbon fuels and digesting carbohydrates both

involve the formation of CO

molecules and the release of

chemical energy. Which type of compounds—hydrocarbons,

carbohydrates, or carbon dioxide—would release the least

amount of energy? Explain how you came to this conclusion.

16. During photosynthesis, green plants take in CO

and com-

bine it with water to make carbohydrates and other com-

pounds. This process is driven by energy from the Sun. When

these food molecules are digested, energy is released. (This

energy can be used to maintain healthy body temperature and

to run other reactions.) Is the energy released via digestion the

same as the energy that absorbed from the Sun? Explain.

17. On a surface excursion during Apollo 14, Alan Shephard hit a

few golf balls. While he originally stated that the balls ﬂew for

miles and miles, he later and more accurately estimated the

distance as 200 to 400 yards. How would the amount of force

needed to drive a ball 300 yards on the Moon differ from that

needed to drive a ball the same distance on Earth? Base your

explanation on the differences between the environments.

18. There are many forces at work in our everyday commute to

work. Identify the typical forces that you’d expect to see

evidence of as you drive an automobile down the street.

19. In the catabolic biochemical pathways in cells, adenosine

triphosphate (ATP) is produced as part of an endothermic

reaction. What is the sign for q in the reaction that represents

the production of ATP? Later, the ATP can be used to help the

cells do work on the surroundings. What is the sign for w in

the reaction that represents that work?

20. In each of the following situations, supply the correct sign for

q and w in the reaction that represents the process.

a. In order to make soft drinks have ﬁzz, or become

carbonated, gaseous CO

at room temperature must be

pumped, under pressure, into an aqueous solution at

lowered temperatures. The system may be considered to be

the resulting carbonated solution.

b. Water placed in a microwave oven can be made to boil

away into steam. Consider the water in its container to be

the system.

Section 5.2 Keeping Track of Energy

Skill Review

21. Suppose a 275-g apple and a 175-g orange are moving with

the same kinetic energy. If the apple is moving at 15 m/s, how

fast is the orange moving? (Maybe we can compare apples

and oranges after all!)

22. Some of the best tennis players may serve a 56.9-g tennis ball

at a velocity of approximately 115 mi/h. What is the kinetic

energy of the tennis ball in such a serve?

23. A molecule of N

moving through a room may have a veloc-

ity of approximately 420 m/s. What is the kinetic energy of a

molecule of N

24. An automobile is moving at a rate of 31 m/s. At this rate it

has a kinetic energy of 436 kJ. What is the mass of this

automobile?

25. A low-fat popcorn snack advertises that one serving provides

90 Cal. One serving is considered to be 34 g. What is the

energy provided in Calories per gram? What is the energy

provided in calories per gram? What is the energy provided

in joules per gram?

26. A different food item advertises that one serving provides

320 Cal. One serving is considered to be 80 g. What is the en-

ergy provided in Calories per gram? What is the energy pro-

vided in calories per gram? What is the energy provided in

joules per gram?

27. How much energy as heat is needed to raise 50.0 g of water

from 23.0°C to 37.0°C?

28. How much energy as heat is required to raise 77.0 g of water

from 18.0°C to 25.0°C?

29. The British thermal unit is used on some heating and cooling

devices. The BTU is deﬁned as the amount of energy as heat

that will raise one pound of water from 58.5°F to 59.5°F.

Express that amount of heat in joules.

30. A popular instant coffee-ﬂavored drink provides 25 food

calories per serving. Express that quantity in calories, joules,

and kilojoules.

Chemical Applications and Practices

31. One of the reactions used in some hand-warmer packets is

the oxidation of iron to form rust. The formation of 1 mole

of rust produces approximately 410 kJ of heat.

a. If this heat were absorbed by 2000.0 g of water at 22.0°C in

a calorimeter, how hot could the water be made?

b. If only 0.10 mole of rust were formed, how hot could

200.0 g of water at 22.0°C become?

32. Ammonium nitrate can sometimes be used in the chemical

cold pack applied to some sports injuries. If 10.0 g of ammo-

nium nitrate were placed in a coffee cup calorimeter that

contained 100.0 g of water and the temperature changed

from 25.0°C to 18.0°C, what was the heat transferred in

kilojoules?

33. Ethanol (C

OH) is being blended with gasoline mixtures

to produce a fuel for automobiles called gasohol. If the

combustion of 10.0 g of ethanol produces 268 kJ of heat,

how much heat will be produced in burning 1 mole of

ethanol?

34. If the combustion of 48.8 g of methane (CH

) produces

855 kJ of heat, how much heat will be produced in burning

1 mole of methane?

204 Chapter 5 Energy

44. A different bomb calorimeter was calibrated by combusting a

2.250-g sample of methanol (CH

O). The calorimeter tem-

perature rose by 0.522°C. What is the heat capacity of this

calorimeter? The molar energy of combustion for methanol

is known to be approximately −890.3 kJ.

45. a. The heat capacity of a certain bomb calorimeter was cali-

brated at 28.9

◦

. When 1.500 g of an unknown sugar was

combusted in the calorimeter, the temperature rose by

2.56°C. What was the energy of combustion of the sugar?

b. What additional information would we need in order to

report the molar heat of combustion for the sugar?

46. A student constructs a crude “coffee cup” calorimeter that

contains 94.1 g of water, at 22.0°C, in a double cup set up

with a thermometer and cork cover. When an 85.8-g piece

of copper at a temperature of 100.0°C was placed in the

calorimeter, the temperature was noted to equilibrate at

28.0°C. The speciﬁc heat of copper is approximately

0.386

g·

◦

a. Calculate the heat gained by just the water.

b. Determine the heat capacity, in J/°C, for the empty

calorimeter.

47. The molar heat of combustion of propane is −2.2 × 10

kJ.

How many grams of propane would have to be combusted to

raise 1.0 kg of water (the amount you might boil to prepare a

tasty macaroni and cheese dinner) from 22.0°C to 100.0°C?

Assume that no heat is absorbed by the container or the air.

48. The speciﬁc heat of water is 4.184 J/g°C. How many grams

of water at 85.0°C would have to be added to raise 1.00 kg of

water from 25.0°C to 50.0°C? Assume that no heat is ab-

sorbed by the container or the air.

49. The heat of combustion for acetylene (C

) is −1300 kJ/

mol. Methane has a heat of combustion of −890 kJ/mol.

Decide, by calculation, which provides more energy as heat

per gram.

50. The heat of combustion for glucose (C

) is

−2803 kJ/mol. Tristearin (C

110

), a typical fat, has a

heat of combustion of −37,760 kJ/mol. Decide, by calcula-

tion, which provides more energy as heat per gram.

51. One way to test a metallic sample to see whether it is made of

gold is to heat it up and place it in a calorimeter to determine

the speciﬁc heat of the sample. The speciﬁc heat of gold

is approximately 0.13

g·K

. What temperature change would

prove that the metal was gold, if 15.0 g of the sample at

99.0°C were placed in a calorimeter initially at 25.0°C? The

heat capacity of the calorimeter is 25.0

52. In the chapter, a bomb calorimeter was discussed that had a

heat capacity of 32.4

◦

. If 0.550 g of benzoic acid (C

)

were combusted in the calorimeter, what would be the

expected change in temperature? (The molar energy of com-

bustion for benzoic acid is −3227 kJ/mol.)

Section 5.3 Speciﬁc Heat Capacity and Heat Capacity

Skill Review

35. Most of the deﬁnitions that you have learned in chemistry

have very speciﬁc terminology. Explain why the deﬁnition of

speciﬁc heat allows you to report the value in either °C or K.

36. In order to clarify the differences among speciﬁc heat, heat

capacity, and molar heat capacity, provide a brief deﬁnition

of each.

37. Which of the following involves the greater amount of heat

transfer?

a. 10.0 kg of water in an automobile cooling system changes

from 45.0°C to 10.0°C as it cools overnight.

b. In preparation for the annual “chili cook-off,” 8.0 L of

water is heated from 22.0°C to 99.0°C.

38. If the heat capacity of a metal coffee pot ﬁlled with water

were known to be 5.87

°C

, what would you calculate to be the

amount of heat transferred from a campﬁre when the tem-

perature of the coffee pot changes from 22.0°C to 95.0°C?

39. Fill in the values missing from the following table.

Chemical Applications and Practices

41. Some research is being done on the production of nickel

compact discs. These disks could store tremendous amounts

of information and last an unusually long time. Like most

metals, nickel has a low speciﬁc heat. If 35.0 g of nickel

absorbs 311 J, it increases in temperature by 20.0°C. What is

the speciﬁc heat of nickel?

42. If 50.0 g of a metal alloy absorbs 471 J, it increases in temper-

ature by 30.0°C. What is the speciﬁc heat of this metal alloy?

43. In order to calibrate a bomb calorimeter, a 2.000-g sample of

benzoic acid (C

COOH; molar mass = 122.12) was com-

busted. The calorimeter temperature rose by 1.978°C. What

is the heat capacity of this calorimeter? The molar energy of

combustion for benzoic acid is known to be approximately

−3227 kJ.

Heat, q Speciﬁc Heat



g·

◦



Mass (g) T (°C)

10.0 joules 4.184 10.0

0.115 10.0 5.0

15.5 joules 42.5 15.0

Heat, q Speciﬁc Heat



g·

◦



Mass (g) T (°C)

4.184 100.0 40.0

450 joules 0.315 15.0

48.5 joules 48.5 15.0

40. Fill in the values missing from the following table.

Chemical Applications and Practices

63. Explain why the values of H and U for reactions such as

the explosion of nitroglycerin, represented here, could vary

signiﬁcantly.

(l) → 12CO

(g) + 10H

O(g) + 6N

(g) + O

(g)

Focus Your Learning

205

62. Of the following, which would not be considered appropriate

equations to represent “standard heat of formation” pro-

cesses? Explain the reasons for your choices.

a. C(s) + O

(g) → CO

(g)

b. H

(g) + O

(g) → H

(g)

(g) + O

(g) → H

O(g)

d. CH

(g) + 2O

(g) → CO

(g) + 2H

O(g)

e. 2H(g) + O(g) → H

O(l)

Section 5.4 Enthalpy

Skill Review

53. Explain how chemical reactions that generate gases can be

thought of as constant-pressure situations but the volume of

the products can be thought of expanding?

54. When the equation U = q

+ w is recast as U = q

−

PV, the sign between the heat and work function changes.

Explain why it is now proper to express the PV with a

negative sign.

55. Without using any symbols or numbers, deﬁne the term

enthalpy as it is applied in chemistry.

56. What is the chief thermodynamic difference among the

expressions known as “heat of reaction,” “standard heat of

reaction,”“standard heat of formation,”and “standard heat of

combustion”?

57. In an exothermic reaction, does the system or the surround-

ings gain energy as heat?

58. What sign would characterize the value of H for the reverse

of an exothermic reaction?

59. What are the conditions being described when a substance is

taken to be in its standard state?

60. What is the physical state of carbon dioxide when it is in

standard state conditions?

61. Of the following, which would not be considered appropriate

equations to represent “standard heat of formation”

processes? Explain the reasons for your choices.

a. C(s) +

(g) → CO(g)

b. N

(g) + 3H

(g) → 2NH

(g)

c. CS

(g) → CS

(l)

d. C(g) + 4H(g) → CH

(g)

e. 4CO

(g) + 5H

O(l) → C

(g) +

(g)

Benzoic acid

64. Consider the reaction for nitroglycerin decomposition

shown in Problem 63. Cite two reasons why the reverse of the

reaction would not qualify as the standard heat of formation

for nitroglycerin.

65. a. Write out the balanced chemical equation for the standard

formation of carbon monoxide.

b. If the value for the standard heat of formation for carbon

monoxide were −110.5 kJ/mol, would you consider the

reaction endothermic or exothermic?

c. What would be the H value for the reverse of the reaction?

66. a. Write out the balanced chemical equation for the standard

formation of gaseous hydrogen peroxide (H

b. If the value for the standard heat of formation for hydro-

gen peroxide were −136.1 kJ/mol, would you consider the

reaction endothermic or exothermic?

c. What would be the H value for the reverse of the reaction?

(gas) (gas) (gas) (gas)

(gas) (gas) (gas)

Nitroglycerin

206 Chapter 5 Energy

ATP

Chemical Applications and Practices

73. The fuel used in many rural settings is propane (C

). Write

out the heat of formation reaction for this important fuel.

Now use the combustion reaction sequence shown in the text

to develop a Hess’s law scheme that would allow you to cal-

culate the H of formation for propane.

74. The sugar arabinose (C

) is obtained from plants with

the polysaccharide gum arabic. In wheat plants, this sugar

helps form important cell wall structures. The heat of forma-

tion reaction for arabinose is

5C(s) + 5H

(g) +

(g) → C

(s)

Without using actual kJ values, show the reactions that you

could use, and the way you would use them, to obtain the

arabinose formation reaction.

75. Given the following reactions and H values:

(s) + 3H

O(g) → B

(g) + 3O

(g) H =+2035 kJ

O(l) → 2H

O(g) H =+88 kJ

(g) +

(g) → H

O(l) H =−286 kJ

2B(s) + 3H

(g) → B

(g) H =+36 kJ

Calculate H for

2B(s) +

(g) → B

(s) H = ?

76. Use the heat of formation reaction data for water and carbon

dioxide and the following reaction for the combustion of

acetylene, employed in high-temperature welding applica-

tions, to ﬁnd the heat of formation of acetylene (C

(g) +

(g) → 2CO

(g) + H

O(l)



H =−1300 kJ/mol

77. a. Ethanol (C

OH) is being added to gasoline to create the

fuel “gasohol.”Using heat of formation data and Hess’s law,

determine the heat of combustion, in kJ/mol, for ethanol.

Combustion of ethanol:

OH(l) + 3O

(g) → 2CO

(g) + 3H

O(l) 

H = ?

Formation of ethanol:

2C (graphite) + 3H

(g) +

(g) → C

OH(l)

H =−278 kJ

b. Using your answer from part a, determine the heat

released when 100.0 g of pure ethanol is combusted.

78. The oxidation of sulfur has many important environmental

connections. Notably, acid rain is formed from sulfur oxides

reacting with moisture in the air. Use the following two reac-

tions to determine the enthalpy change when sulfur dioxide

reacts with oxygen to form sulfur trioxide.

a. S(s) + O

(g) → SO

(g) H° =−296.8 kJ

b. 2S(s) + 3O

(g) → 2SO

(g) H° =−791.4 kJ

2SO

(g) + O

(g) → 2SO(g) H° = ?

79. The octane rating on gasoline is a method of comparing fuel

energy values. What is the value for the heat of combustion of

octane? (

H°(C

) =−249.9 kJ/mol)

(l) +

(g) → 8CO

(g) + 9H

O(l)

80. A typical componentof gasoline is pentane (C

).The stan-

dard enthalpy of combustion of pentane is −3537 kJ/mol.

Write the combustion reaction for pentane, and determine

the standard enthalpy of formation for pentane.

70. Glucose (C

) is used by our bodies as an energy source.

Write out the standard heat of formation reaction for this

important molecule.

Section 5.5 Hess’s Law

Skill Review

71. The H value for the following reaction is −1012 kJ.

2NH

(g) + 3N

O(g) → 4N

(g) + 3H

O(l)

a. What is the value of H for the reverse of the reaction?

b. What is the value of H for 1 mole of NH

reacting?

c. What is the value of H of 4 moles of NH

reacting?

72. The H value for the following reaction is +284.6 kJ.

(g) → 2O

(g)

a. What is the value of H for the reverse of the reaction?

b. What is the value of H for 1 mole of O

reacting?

c. What is the value of H of 4 moles of O

reacting?

67. The hydrocarbon fuel butane (C

) is used in small

portable lighters. Write the reaction for the standard heat of

combustion for this reaction. (Recall that the two products of

hydrocarbon combustion are carbon dioxide and water.)

68. Pentane (C

) can be blended with other hydrocarbons

and additives to form gasoline. Write the reaction that de-

picts the standard heat of formation of pentane.

69. The primary energy molecule in cells is ATP (adenosine

triphosphate, C

). Write out the standard heat

of formation reaction for this important biomolecule.

Focus Your Learning 207

81. Calcium metal will react in water to form calcium hydroxide,

Ca(OH)

.Use the thermochemical data that follow and Hess’s

law to calculate the value of H° for the reaction

Ca(s) + 2H

O(l) → Ca(OH)

(s) + H

(g)

a. H

(g) +

(g) → H

O(l) H° =−286 kJ

b. CaO(s) + H

O(l) → Ca(OH)

(s) H° =−64 kJ

c. Ca(s) +

(g) → CaO(s) H° =−635 kJ

82. Ozone is reduced by hydrogen to produce water. Use the

thermochemical data that follow and Hess’s law to calculate

the value of H° for the reaction

(g) + O

(g) → 3H

O(g)

a. H

(g) +

(g) → H

O(l) H° =−286 kJ

b. 2H

(g) + O

(g) → 2H

O(g) H° =−483.6 kJ

c. 3O

(g) → 2O

(g) H° =+284.6 kJ

Section 5.6 Energy Choices

Chemical Applications and Practices

83. Most of the world’s energy consumption for power is based

on fossil fuels. This source, however, is considered nonre-

newable. Look back at the list of renewable energy sources

and cite one advantage, in addition to this renewability,

that each would have over the petroleum-based fuels used

today.

84. Assuming that the fossil fuels are completely depleted at

some point in the future, the world will need to depend on

renewable sources of energy. Look back at the list of renew-

able energy sources and describe one disadvantage of using

each source as the sole source of energy for a country.

Comprehensive Problems

85. Individual atoms and molecules are so small that they have

very low values of kinetic energy. However, given their mass

and velocity, it is possible to calculate the value. What is the

kinetic energy of an oxygen molecule (O

) in air that you are

breathing if its velocity is 460 m/s? Would you expect a

nitrogen molecule (N

) moving at the same speed to have

more or less kinetic energy than the oxygen molecule?

Explain.

86. Distinguish between the two terms in each of the following

pairs:

a. Heat and temperature

b. System and surroundings

c. Exothermic and endothermic

d. q and U

87. What is the role of activation energy in a chemical process? Is

it possible for the activation energy of a reaction to be greater

than the overall change in energy for a chemical reaction?

Explain.

88. The average temperature of a healthy human is approxi-

mately 37°C. The average room temperature may be about

25°C. Explain how we are able to keep our body temperature

so much higher than that of our environment.

89. A 400-mL glass beaker contains 250 g of water at room

temperature. As several NaOH pellets are placed in the water

and begin to dissolve, you notice a warming sensation in the

hand in which you are holding the beaker. Answer the fol-

lowing questions.

a. If the NaOH and the water make up the system, would you

consider the process endothermic or exothermic? Explain.

b. Is the beaker part of the system or part of the

surroundings?

c. During this process, is energy ﬂowing into or out of the

system?

d. During this process, has the kinetic energy of the water

molecules been raised or lowered?

e. What work is being done during this process?

90. A serving of Italian rice—risotto—provides 150 food Calo-

ries. What would this value be in kilojoules? Assume that this

quantity of energy would be available to do the work of lift-

ing 2.5-kg chemistry textbooks from the ﬂoor to a height of

1.5 m. How many such chemistry textbooks could, theoreti-

cally, be lifted through that height? (And later, of course, after

that work, brought to class.)

91. Express the energy content of a 2.00-oz. candy bar that con-

tains 247 Calories in Calories per gram, in joules per gram,

and in kilojoules per gram. If this energy were efﬁciently

used to provide the kinetic energy to move a 1.5-pound

chemistry book, how fast, in meters per second, could the

book move?

92. One of the main reasons for eating is to obtain a supply

of energy. A low-fat apple mufﬁn may provide 170 food

Calories for each 50-g mufﬁn. How many mufﬁns (or frac-

tions of mufﬁns) are needed to provide 1 joule?

93. a. Suppose you are heating water (225 g) in a mug that you

have placed in a microwave oven. As you wait to add the in-

stant hot chocolate, please calculate, from the following

data, the amount of energy as heat that the water has ab-

sorbed: The original water temperature was 15.0°C. When

you remove the mug of hot water, you ﬁnd that the tem-

perature has risen to 98°C.

b. What additional information would you need in order to

determine the heat absorbed by the mug?

94. You have just removed a hot cheese pizza from the oven, and

all of the ingredients are presumably at the same tempera-

ture. Without waiting for it to cool, you take a bite of the

pizza. As you bite the pizza, the bread is hot on your tongue

but does not burn. However, as you continue to bite, pizza

sauce (mostly tomatoes and water) squeezes out and burns

the roof of your mouth. Which has the higher speciﬁc heat,

the bread or the sauce? Explain the basis of your answer using

q = m × c × T.

95. One way to determine the heat capacity of a constant-vol-

ume calorimeter is to burn measured amounts of pure car-

bon in the presence of oxygen gas to form carbon dioxide.

From the following experimental data, determine the heat

capacity of the calorimeter. A 0.200-g sample of carbon,

when completely combusted, raised the temperature of the

water and the contents of the entire calorimeter from 24.0°C

to 25.5°C. It is known that under these conditions, the heat

released from the complete combustion of 1 mole of carbon

is 392 kJ.