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

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MEDIA ENHANCED EDITION

Houghton Mifflin Company

Boston New York

Chemistry

The Practical Science

Paul Kelter

Northern Illinois University

Michael Mosher

University of Nebraska at Kearney

Andrew Scott

Perth College, UHI Millennium Institute

Contributing Writer

Charles William McLaughlin

University of Nebraska at Lincoln

Publisher: Charles Hartford

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Credits appear on pages A49–A52, which are considered an extension of the copyright page.

No part of this work may be reproduced or transmitted in any form or by any means, electronic or

mechanical, including photocopying and recording, or by any information storage or retrieval system

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Library of Congress Control Number: 2007937002

ISBN-10: 0-547-05393-2

ISBN-13: 978-0-547-05393-6

123456789–CRK–11 10 09 08 07

To Barb, Kristie, and Margaret,

who are with us through it all, and to Seth, Aaron, Jamie, David, and

Alison, for whom we do these things

To Jim Carr,

master craftsman of the finest teachers

Develop effective

problem-solving skills

First Thoughts encourages

students to think about a

strategy before attempting

the solution.

Solution demonstrates the

detailed solution to the

problem and effectively

uses color to highlight key

information.

Further Insights continues

the discussion of the

problem, illustrating

applications of practical

importance and linking

multiple concepts

throughout the book.

Practice includes addi-

tional practice problems

for students to try on

their own, with answers

in the appendix.

EXERCISE 5.3 The First Law of Thermodynamics

Calculate the change in the energy of a system if 51.8 J of work is done by the system

with an associated heat loss of 12.3 J.

First Thoughts

We must pay particular attention to the sign conventions for heat and work in this

problem. In this case, work is done by the system. Is this work positive or negative?

A useful system to keep in mind is you! That is, when you do work—by running,

dancing, or even moving your textbooks from one class to the next—you are using

energy. After the process of moving your body or your books, you have less energy

than you had before, so the work has a “−” sign. Similarly, the 51.8 J of work done

by the system means that its energy change, w

system

,is−51.8 J. The energy loss as

heat by the system (such as that accompanying your run as your body tries to stay

cool) also has a “−” sign, so q

system

=−12.3 J.

Solution

From the standpoint of the system,

U = q + w

U =

−

12.3 J + (−51.8 J) =

−

64.1 J

Further Insights

We want to reinforce the point that there is no such thing as heat or work within a

system. In other words, a system does not contain heat or work. Rather, heat and work

exist only as types of energy transfer between the system and the surroundings. Work

is done by or on a system to move it through a distance. A system can transfer 64.1 J

of energy as heat and work. It does not contain 64.1 J of heat and work.

PRACTICE 5.3

Calculate the change in the energy of a system if 84.7 J of work is done on the sys-

tem, with an associated loss of energy as heat of 39.9 J.

See Problems 11–14, 19, 20, and 89.

The problem-solving approach is more about reason than rote. Before jumping

headfirst into an Exercise, students are asked to consider First Thoughts. These

remarks often ask students to apply their real-world experience to generate

ballpark answers. They also call on prior course content to help students see the

big picture and develop a plan of attack. A detailed solution is then followed by

Further Insights, which compel students to not only ask themselves if their an-

swer makes sense but to put the entire exercise in context and make connections

to biology and numerous other disciplines. In short, students are asked to think

like a scientist.—Ed O’Sullivan, Parkland College

This way of doing things gives students real insights into how “real chemists or

scientists” think.—Milton Johnston, University of South Florida

This approach to problem solving emphasizes problem strategies, and

then encourages students to look beyond the problem.

Engage students through

integrated applications

A space shuttle launch is an awe-inspiring

demonstration of the ability of energy to trans-

port humans and material away from Earth’s surface

with an eye toward planetary exploration. The energy to

launch the craft and its crew comes from the explosive vio-

lence of chemical reactions. When these reactions are used in a

carefully controlled way, they can lift the massive shuttle (which

weighs in at a robust 2,000,000 kg), the booster rockets, and the fuel tank

and propel the ship into orbit hundreds of miles above Earth in a very brief

10-minute ride.

Two chemical reactions, indicated in Figure 5.1, power the launch of the space

shuttle. The combustion of hydrogen gas (the combination of hydrogen and oxygen

to form water) takes place in the main engines. At the same time, the solid rocket

boosters attached to the sides of the shuttle release a host of products resulting from

the oxidation of aluminum by ammonium perchlorate, though we show only the

primary equation here.

Main engines: 2H

(g) + O

(g) →2H

O(g)

Boosters: 10Al(s) + 6NH

ClO

(s) → 5Al

(s) + 6HCl(g) + 3N

(g) + 9H

O(g)

As the shuttle sits on the launch pad, it is hard to believe that the en-

ergy that will launch it with such a spectacular display of chemical mus-

cle is quietly present within the main fuel tank and the solid fuel of the

booster rockets. Chemicals can store huge amounts of energy in this

Application

HEMICAL ENCOUNTERS:

Setting the Stage with

the Space Shuttle

Interwoven throughout the chapter,

chemistry concepts are explained

through the use of real-world examples.

Application icons throughout the text

show at a glance where chemical

concepts are applied.

Applications are also found

How Do We Know? and

NanoWorld / MacroWorld

features, chapter openers,

photos, and end-of-chapter

exercises. Students will see how

chemistry connects to their lives,

their future careers, and their

world.

4. Diagram, using circles for atoms, a crystal of KCl versus the

same crystal of KCl dissolved in water.

Chemical Applications and Practices

5. Earth’s oceans contain tons of dissolved sodium chloride.Yet,

when a ship develops an oil leak, almost none of the oil dis-

solves in the ocean. Explain this phenomenon.

6. When an ion dissolves, it is surrounded by a hydration

sphere. If water molecules surrounded the ion so that the hy-

drogen portion of the water was closer to the ion, would the

ion most likely be a cation or an anion?

7. Pure water does not conduct an electric current. However,

aqueous solutions of some compounds do form solutions

that conduct electricity. Explain why the presence of some

solutes converts nonconducting water into a conducting

solution.

8. Glycerin can be produced as a by-product in soap making.

The compound dissolves so easily in water that it absorbs

water from the air. This latter characteristic is why glycerin

is often found in many skin lotions. As glycerin absorbs

the water, the skin can be kept moist. Glycerin’s structure is

shown below. What aspects of glycerin’s structure contribute

most to its ease of dissolving in water?

9. A conductivity-testing apparatus, such as the one shown in

this chapter, possesses a light bulb whose brightness is related

to how much current is ﬂowing through it (and also through

the solution.) A small, but measurable, amount of current

must be present before the bulb becomes visibly brighter.

What effect would this characteristic of the apparatus have

12. Which of the following would best represen

water?

CHH

cerin

(a) (b)

OCl

–

(a) (b)

“Smile, you’re on candid camera!” Everyone hopes never

to be ambushed with those words. In the nineteenth cen-

tury, when photography was in its infancy, candid pho-

tographs were not possible. In order to create a visual

memory of an event, you had to visit the photographer

and sit for a picture.“Sitting” meant holding a pose for

up to 30 seconds. Any movement would show as a blur

in the ﬁnal picture. However, the idea that an inexpen-

sive, yet realistic, picture could be created in such a short

time brought people to the photographer’s studio in

droves.

Scientists in the early 1800s were intrigued by the

idea of painting portraits without the artist’s pen. In the

summer of 1827, Joseph Nicéphore Niépce, using mater-

ial that hardened on exposure to light exposed ﬁlm for

eight hours in order to get a picture. Because no one

wanted to sit still that long, most of his pictures were

landscapes. In 1839, however, two other photographic

processes were developed, the daguerreotype (perfected

by Louis Daguerre) and the calotype (Henry Fox Talbot).

Although these techniques allowed pictures to be created

with much shorter exposures, they still produced a

black-and-white image (see Figure 9.30). This didn’t

matter much to the general public, and photographic

portrait studios seemed to pop up on every corner across

the globe. Despite some marginal success in producing

faint color images during the 1850s and 1890s, the ad-

vent of the color photograph as an inexpensive, repro-

ducible and stable process didn’t occur until the 1920s.

Researchers at Kodak Research Laboratories in 1935

ﬁnally hit upon the process that would bring the color

photograph to the world. Kodachrome®, their color

photography process, introduces color to a picture

by separating the three primary colors into speciﬁc

emulsion layers. How does it work? The emulsion layers

contain a compound that reacts when light strikes it. A

quick look at the chemistry behind the black-and-white

photograph will give us some insight.

In black-and-white photography, a piece of transpar-

ent plastic is coated with silver bromide crystals (see Fig-

ure 9.31) and a sensitizer. The plastic sheet is then placed

in a camera, and when the shutter is opened, the silver

bromide absorbs photons of light. In the presence of the

sensitizer, the silver cation is reduced by the sensitizer

(gains an electron) and becomes silver metal:

(s) + sensitizer n Ag°(s) + sensitizer

Then the ﬁlm is wound back into its container and

sent to be developed. The technician rinses the plastic

ﬁlm to remove unreacted silver bromide. The silver

metal remains in place on the plastic. The result is an in-

verted image called a negative. Next the technician shines

light through the negative onto another silver bromide/

sensitizer-coated support (typically paper this time) and

rinses off the unreacted silver bromide. Violà! We have a

positive image called a photograph.

Why is the ﬁlm coated with silver bromide instead of

silver chloride? We can answer this by examining the

wavelength of light that is absorbed; see Figure 9.32.Com-

parison of the two shows that silver bromide absorbs a

large amount of the visible light that enters our camera.

To get the best picture, we want to absorb as much light

as we can.

NanoWorld / MacroWorld

Big effects of the very small:

Color photography and the tricolor process

FIGURE 9.30

A daguerreotype of Michael Faraday

(see Chapter 19) taken sometime

between 1844 and 1860 at Mathew

Brady’s studio in New York City.

Because the process required the

subject to remain completely

motionless for up to 10 seconds,

most daguerreotypes were taken

in well-lit studios. Even though

the early daguerreotypes cost one

month’s wages for the average

person, they were much cheaper

than a painted portrait. This made

them an instant success.

FIGURE 9.31

A magniﬁed image of silver bromide crystals used in the photographic

industry. Note that the crystals are mostly hexagonal in shape and ﬂat

so that more surface area is exposed to incoming photons.

vii

Dilution

How do municipal water treatment plant workers prepare water so that it contains

a ﬂuoride ion concentration of about 1 part per million?

They use a concentrated

source of ﬂuorine, hydroﬂuorosilicic acid, H

SiF

(“HFSA”), which they then di-

lute in the drinking water. HFSA reacts with water in a fairly complex way that re-

leases ﬂuoride ion into the water. In Ireland, a company in the town of New Ross,

Application

Ask questions

Inquiry-based learning is integral to this book. Questions are asked throughout

the narrative to help students make the connection between what they are

reading, concepts throughout the text, and their everyday lives.

Highlighted questions

appear throughout the

chapter and prompt

students to learn how to

raise meaningful questions

about topics.

How do we know? essays

ask questions to probe

the use of key chemical

concepts in medicine,

research, and other

applications.

Students are also given

many opportunities to

reflect on what they have

learned throughout the

book.

Here’s What We

Know So Far in-chapter

summaries review

complex topics.

The moons of the outer planets make

most elegant chemical laboratories,

because conditions on these worlds are so very dif-

ferent from those here at home. When we look at

photos of Jupiter’s moons, Europa and Ganymede

(Figure 11.17), we see some vexing geological fea-

tures. How do we know what caused these features?

We can’t visit the moons, at least not directly. How-

ever, our space probes can gather various types of

electromagnetic radiation—including light, ultra-

violet radiation, and X-rays—that give us data

from which to draw conclusions.

The probes’ data seem to show that deep within

the surface of these moons, there are layers of ice,

each in a different phase, mixed with rocks. The phase

diagram of water that we showed as Figure 11.18

is inadequate to account for the low temperatures

and massive pressures found within these moons.

A more comprehensive phase diagram for water,

emphasizing its solid phases, is given as Figure 11.20.

The ice that is made in our kitchen freezer or found

in an ice storm—what we might call “normal ice”—

is technically called Ice Ih. Its structure is shown in

Figure 11.21. There are 11 other forms

of ice that have been made in the laboratory or

simulated by computer.

The phase diagram shows that as the

temperature and pressure inside the moons

change, different types of ice are formed.

Each ice has its own density. As layers

form versions of ice of different

densities, they form the geological

features, such as cracking and

How do we know?

The moons of Jupiter

0 100 200 300 400 500 600 700 800

Solid

IhIc

III

VII

VIII

Liquid

Vapor

Temperature (K)

Pressure (Pa)

FIGURE 11.20

This is an expanded phase diagram for water, taking into account

the phase changes among the various solid (ice) forms. These forms

of ice are thought to occur in layers beneath the surface of some

of the moons of the outer planets.

FIGURE 5.11

An endothermic process. Some chemical

reactions require the transfer of energy

from the surroundings to the system. This

is an endothermic process.

HERE’S WHAT WE KNOW SO FAR

■

There are three natural forces: the gravitational, electroweak, and strong

nuclear forces.

■

Potential energy is the energy stored within a system.

■

Kinetic energy is the energy associated with motion.

■

The law of conservation of energy states that energy can be neither created nor

destroyed. Instead, it just moves between the system and the surroundings

and can be transformed from one form to another.

■

Work and heat are two ways in which energy can move between the system

and the surroundings.

■

The ﬂow of energy as heat from the system to the surroundings involves an

exothermic process. An endothermic process involves the ﬂow of energy as

heat from the surroundings to the system.

■

Changes in the energy of a reaction can be studied by examining a reaction

proﬁle diagram.

Cold packs use endothermic reactions.

viii

Dynamic, contemporary

art program

Let’s look at these concentration units in a slightly different way.As an example,

the EPA maximum level for hexachlorobenzene (C

), a pesticide used onwheat

in the United States until 1965, is 1 ppb. This means that the maximum allowable

level in drinking water should be 1 gram of C

per billion grams of solution.

1ppbC

1gC

g solution

If we use

1 g water

1 mL water

as the density of water and make the key assumption that the

solution is so dilute that its density is about equal to that of water, then

Density of very dilute aqueous solutions

1 g solution

1 mL solution

This means that we can express the concentration of hexachlorobenzene as the

number of grams of C

per liter of solution:

1gC

g solution

1 g solution

1 mL solution

mL solution

1 L solution

−6

L solution

1 µgC

L solution

↑↑ ↑

(one ppb C

) (density of the solution) (convert mL to L)

The bottom line is that in dilute aqueous solutions, we can express parts per

billion as

136 Chapter 4 Solution Stoichiometry and Types of Reactions

Hexachlorobenzene

(a) (b) (c)

FIGURE 4.10

To visualize the idea of parts per million,

per billion, and per trillion, consider that

one drop of ink could be placed in these

quantities of water. (a) Placing that drop

of ink in a 12-gallon bucket results in

1 ppm. (b) Placing it in a tanker truck

results in 1 ppb. (c) Placing it in a

12-million-gallon reservoir results in

1 ppt.

This contemporary art and photo program appeals to tech-savvy students who

expect exciting and visually appealing graphics regardless of medium.

11.2 A Closer Look at Intermolecular Forces 445

Water boils and creates steam.

a mole of liquid methane, compared to 1650 kJ to break the four C—H bonds in

a mole of methane! We can extend this information to make the more general

statement that intermolecular interactions are much weaker than intramolecular

interactions. Less energy is required to boil a liquid than to break it into its

component elements.

11 2 ACl L k I l l F

Five molecules of liquid water interact

with each other. Note the positions of

the hydrogen atoms. They are being

shared with neighboring oxygen atoms.

Molecular-level illustrations

of key concepts help students

connect nanoscale activity to

macroscale phenomena.

In addition,

electrostatic

potential maps use vibrant

color to show how

electron density changes

across a molecule.

128 Chapter 4 Solution Stoichiometry and Types of Reactions

Ethanol

An electrostatic potential map of glucose and

ethanol. The electrostatic potential is plotted

on the surface of a computer-generated

model of the molecules. Note how the electron

density is distributed in each molecule, and

compare these models to that of water

(Figure 4.3).

D-glucose

A wealth of end-of-chapter

problems test students’

understanding of key

concepts and problem-

solving skills. They include

Skill Review and Chemical

Applications and Practices

organized by chapter

sections and in matched

pairs.

Comprehensive Problems

build on the students’

mastery of concepts in

the chapter and across

multiple chapters in the

textbook.

Thinking Beyond the

Calculation provides

rigorous, cumulative, and

conceptual problems based

on multiple concepts.

Focus Your Learning

The answers to the odd-numbered problems appear at the back of

the book.

Section 4.1 Water—A Most Versatile Solvent

Skill Review

1. Explain how water molecules can dissolve both cations and

anions.

2. Why doesn’t pure water conduct an electric current?

3. Explain what is meant by the term hydration sphere?

4. Diagram, using circles for atoms, a crystal of KCl versus the

same crystal of KCl dissolved in water.

Section 4.2 The Concentration of Solutions

Skill Review

11. Which of the following would best represent MgCl

dissolved

in water?

(a) (b)

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

)pj

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,

a. If the value of H for the reaction is −453 kJ, what is the

value of H for the reverse of the reaction?

b. What is the value of H for this reaction if 10.0 g of

phosphoric acid is produced?

c. Is this reaction endothermic or exothermic?

d. If 1.50 g of P

(s) and 2.50 mL of water were mixed,how

many grams of phosphoric acid would result?

e. What is the enthalpy change for the process outlined in

part d?

f. If 10.0 g of P

were mixed with 1.00 kg of water at

25.0°C, what would be the ﬁnal temperature of the water?

22.7°C

152.06 g =

Mass of

calorimeter

Water

added

Metal block

at 96.3°C

234.95 g =

Mass of

calorimeter

and water

257.88 g =

Mass of

calorimeter,

water, and

metal block

24.1°C

Phosphoric acid

96. A student’s coffee cup calorimeter, including the water it

contains, has been calibrated in a manner similar to that de-

scribed in Problem 46. The heat capacity was found to be

55.5

◦

. If a 65.8-g sample of an unknown metal, at 100.0°C,

was placed in the calorimeter initially at 25°C, and an equi-

librium temperature of 29.1°C was reached, what is the

speciﬁc heat of the metal?

97. The foods we eat provide fuel to keep us alive. Burning a

0.500-g sample of vegetable oil provides enough heat to raise

the temperature of a calorimeter by 2.5 K. Assuming the heat

capacity of the calorimeter to be 7.5

, determine the heat of

combustion for 1 g of the oil.

98. A student performs the experiment shown graphically here.

What is the speciﬁc heat of the block of metal used in the ex-

periment? (Assume that the heat capacity of the empty

calorimeter is 7.5

◦

Thinking Beyond the Calculation

99. Phosphoric acid is used in many soft drinks to add tartness.

This acid can be prepared through the following reaction:

(s) + 6H

O(l) → 4H

(aq)

Comprehensive

end-of-chapter materials