Chemistry 4ed 2005 olmsted williams

About the AuthorsPreface 1 The Science of Chemistry 2 The Atomic Nature of Matter 3 The Composition of Molecules 4 Chemical Reactions and Stoichiometry 5 The Behavior of Gases 6 Energy a

Copyright © 2005 John Wiley & Sons, Inc All rights reserved.

John Wiley & Sons, Inc

2005

About the Authors

Preface

1 The Science of Chemistry

2 The Atomic Nature of Matter

3 The Composition of Molecules

4 Chemical Reactions and Stoichiometry

5 The Behavior of Gases

6 Energy and Its Conservation

7 Atoms and Light

8 Atomic Energies and Periodicity

9 Fundamentals of Chemical Bonding

10 Theories of Chemical Bonding

11 Effects of Intermolecular Forces

12 Properties of Solutions

13 Macromolecules

14 Spontaneity of Chemical Processes

15 Kinetics: Mechanisms and Rates of Reactions

16 Principles of Chemical Equilibrium

17 Aqueous Acid–Base Equilibria

19 Electron Transfer Reactions

20 The Transition Metals

21 The Main Group Elements

22 Nuclear Chemistry and Radiochemistry

A Scientific Notation

B Quantitative Observations

C Ionization Energies and Electron Affinities of the First 36 Elements

D Standard Thermodynamic Functions

E Equilibrium Constants

F Standard Reduction Potentials,

Solutions to Odd-Numbered Problems

Answers to Problems

1 The Science of Chemistry

Matter Is Molecular

A view of the Earth from space shows that our planet is an integrated whole At the same time, we know that the Earth is mind-boggling in its diversity Nevertheless, the stunning complexity of our world can be described

using a small set of chemical principles These fundamental aspects of chemistry are the subject of this book

The entire universe is made up of matter, from the vast reaches of the galaxies to a simple glass of water As we

describe in the coming chapters, matter is composed of tiny particles called “atoms.” On Earth there are around

100 different kinds of atoms, each kind with its own unique combination of properties The complexity of our

world arises from the unlimited number of ways that atoms can combine to form different molecules The

principles of modern chemistry are organized around the molecular nature of matter Our book presents this

perspective while at the same time emphasizing the quantitative aspects of chemistry

A drop of water contains an unimaginable number of molecules, as our molecular inset shows Water is essential

to life as we know it The simple yet unusual fact that solid water (ice) floats atop liquid water allows life to exist

on our planet Just as important is the fact that water dissolves an immense range of chemical compounds: Water

is the solvent of life In fact, water is so important to our perspective of life that the search for water is a key

feature of our quest to discover life in other quarters of the galaxy The inset photo of the surface of Mars, for

example, shows no sign of water at present, but some erosional features appear to have been caused by flowing water in the past

chemistry being the same throughout the universe Moreover, from research on stars, chemists have learned that the various kinds of atoms probably form during stellar evolution and are dispersed throughout the universe by supernova explosions.

Chemists are interested in a huge range of problems, extending from the galactic scale to what takes place

between individual atoms and molecules Here are some of the practical problems identified by a 2003 report

from the National Research Council, “Beyond the Molecular Frontier: Challenges for Chemistry and Chemical Engineering.”

• Develop new materials that will protect citizens against terrorism.

• Develop medicines and therapies that can cure currently untreatable diseases.

• Develop unlimited and inexpensive energy to pave the way to a truly sustainable future.

• Revolutionize the design of chemical processes to make them safe and environmentally benign.

• Understand the complex chemistry of the Earth, to design policies that will prevent environmental

Matter is anything that possesses mass and occupies space.

How Chemistry Advances

Chemists learn about chemical properties by performing experiments They organize information about chemical properties using general principles and theories The periodic table, for example, organizes the elements according

to chemical properties Chemists use general principles and theories to make predictions about yet-unknown

substances These predictions generate experiments whose results may extend the scope of the principles and

theories

Chemical research is driven by many goals, and it progresses in many different ways The essential traits of a good researcher are curiosity, creativity, flexibility, and dedication Some chemical advances come from a direct assault

on a known problem A classic example is the development of the Hall process for refining aluminum from its

ores, which we describe in Chapter 21 As a contemporary example, many scientists around the world are working

at a feverish pace to develop a vaccine against the AIDS virus

Chemistry also advances when an imaginative researcher recognizes the potential of a lucky accident Synthetic dye-making, the first major chemical industry, arose from one such lucky accident While searching for a way to synthesize quinine, a drug for the prevention of malaria, English chemist William Perkin accidentally made a

beautiful reddish-violet dye, which he called mauveine Perkin had the imagination to realize the commercial

potential of the new substance and switched his research interests from drugs to dyes His insight made him rich and famous

Mauveine is a vivid red-violet dye

Chemical advances frequently are driven by technology The discovery that atoms have inner structure was an

outgrowth of the technology for working with radioactive materials In Chapter 2 we describe a famous

experiment in which the structure of atoms was studied by bombarding a thin gold foil with subatomic particles A contemporary example is the use of lasers to study the details of chemical reactions We introduce these ideas in Chapters 7 and 8

Methods of Science

However a new chemical discovery arises, an essential component of science is to explain that discovery on the

basis of general principles When a new general principle is posed, it is termed a hypothesis A hypothesis is

tentative until it can be confirmed in two ways First, additional observations must be consistent with the

hypothesis; second, the hypothesis must predict new results that can be confirmed by experiments If a hypothesis

meets these tests, it is promoted to the status of a theory A theory is a unifying principle that explains a collection

of facts

Experimental observations are at the heart of chemical research Many experiments are designed specifically to answer some particular chemical question Often, the results of these experiments are unexpected and lead to new hypotheses New hypotheses, in turn, suggest additional experiments The Chemistry and Life Box describes how the hypothesis of extraterrestrial life can be tested

A typical example of the interaction between hypothesis and experiment is the story of the work that resulted in worldwide concern over the depletion of the ozone layer in the stratosphere These studies led to the awarding of the 1995 Nobel Prize for Chemistry to Paul Crutzen, Mario Molina, and F Sherwood Rowland Figure 1-1

provides a schematic view of how this prize-winning research advanced It began in 1971 when experiments

revealed that chlorofluorocarbons, or CFCs, had appeared in the Earth's atmosphere At the time, these CFCs were

Figure 1-1

This flow chart illustrates how the scientific process led to worldwide concern over the effect of

chlorofluorocarbons on the ozone layer

Meanwhile, Crutzen had done experiments showing that ozone in the upper atmosphere can be destroyed easily by reactions with nitrogen oxides This work demonstrated that the ozone layer is in a delicate balance that could be disturbed significantly by changes in atmospheric composition In 1974, Molina and Rowland combined Crutzen's experimental work with their own theoretical analysis and published a prediction (hypothesis) that CFCs pose a serious threat to the ozone layer

Following this interplay between observations and theory, many atmospheric scientists began studying chemical reactions of ozone in the upper atmosphere Chemists duplicated atmospheric conditions in the laboratory and

measured how fast various chemical reactions occur The results of these experiments were used to create

theoretical models of the upper atmosphere and predict how the ozone concentration would change as CFCs were

introduced Meanwhile, atmospheric scientists carried out measurements showing that ozone was being depleted

in the upper atmosphere at a rate even faster than had been predicted

Today, scientists realize that the chemistry of the upper atmosphere is quite complex In addition to gaseous

molecules, solid particles such as tiny ice crystals play important roles in the chemistry that affects ozone The

original hypothesis of Rowland and Molina, that CFCs reach the upper atmosphere and deplete the ozone layer, has been fully confirmed Exactly how this occurs, what other chemicals are involved, and how this process might

be controlled, are still under intense study by chemists and other scientists, leading to yet more hypotheses and

experiments

The story of research into the depletion of atmospheric ozone is just one example of how scientific understanding and theories develop The fundamental theories and laws of chemistry that we present in this text all went through similar intense scrutiny and study

Chemistry and Life: Is There Life on Other Planets?

Speculation about life on other planets probably began when humans discovered that the Earth is not

unique We know that several other planets of the solar system bear at least some resemblance to our own Why, then, should there not be life on Mars, or Venus, or perhaps on undiscovered Earthlike planets

orbiting some other star?

How can scientists collect experimental evidence about possible life on another planet? Sending astronauts

to see for themselves is impractical at our current level of technology Nevertheless, it is possible to search

for life on other worlds without sending humans into space In the late 1970s, NASA's Viking spacecraft

lander collected a sample of dirt from Mars, the planet in our solar system most like Earth The sample

showed no signs of life Nevertheless, speculation continues about Martian life

The photo below, taken by the Viking spacecraft, shows that the surface of Mars has been eroded,

apparently by liquid water More recent photos transmitted by Spirit and Opportunity convince scientists

that this was the case Apparently, Mars was once much warmer than it is today Planetary scientists

speculate that at one time the atmosphere of Mars may have contained large amounts of carbon dioxide,

setting up a “greenhouse” effect that made the surface of that planet warmer and wetter Might there, then, have been life on Mars at some earlier time? Molecular structures found in meteorites thought to come from Mars have been interpreted to show that there was once life there, but these results are controversial

Indirect evidence can be collected without actually visiting a planet Recent photographs taken from flyby

of ice, which suggests that there may be liquid water under the ice, warmed by tidal forces generated by

Jupiter's huge mass and strong gravity Because life seems to require the presence of water, this observation indicates that there could be life on Europa

Conditions on other planets seem too hostile for life as we know it, but recent discoveries on our own planet indicate that life is much more robust than was once thought Deep-sea explorers have discovered

flourishing life around hydrothermal vents Whereas life on the surface of the Earth relies on sunlight and photosynthesis for energy, these deep-sea life forms exploit energy-rich compounds spewed forth by

volcanic vents The warm waters around hydrothermal vents (photo below) teem with bacteria, which in

turn support higher life forms such as worms and crustaceans This terrestrial life thrives in an environment similar to one that might exist on Europa, reinforcing speculation that this moon of Jupiter could support

some forms of life

Outside our own solar system, might there be planetary environments where life flourishes? In recent years, astronomers have discovered planets orbiting stars other than our own Whether or not these planets support life is still impossible to say Nevertheless, the more we discover about the variety of the universe, the more likely it becomes that we are not alone

Section Exercises

c01-0001.List four ways that chemistry applies to cooking

c01-0002.Describe how chemistry applies to the automobile industry.

c01-0003.A planetary scientist announces a theory predicting that substances on Venus react differently than

they do on Earth Write a paragraph that describes ways to test this theory

Atoms, Molecules, and Compounds

Every substance has physical properties that we can measure and describe, including shape, color, and texture For example, the iron girder shown in Figure 1-2 has a lustrous silvery color and a smooth texture Substances also have chemical properties, such as their flammability Physical and chemical properties that can be observed with the eye

are called macroscopic The underlying structure of a chemical substance, which is called microscopic, can be

explored using magnifying devices The magnified view of iron shown in Figure 1-2 reveals fissures and pits that are not visible in the macroscopic view Still further magnification eventually reveals the building blocks of matter This,

the molecular or atomic view, is an essential part of every chemist's thinking.

Figure 1-2

Iron appears different at the macroscopic, microscopic, and atomic levels

Atoms

The fundamental unit of a chemical substance is called an atom The word is derived from the Greek atomos,

meaning “uncuttable.” An atom is the smallest possible particle of a substance

Atoms are extremely small Measurements show that the diameter of a single carbon atom is approximately 0.000

000 0003 meters (about 0.000 000 001 feet) To give you some idea of just how small that is, a sample of carbon the size of the period at the end of this sentence contains more atoms than the number of stars in the Milky Way Any sample of matter large enough for us to see or feel contains an unfathomable number of atoms

Molecules

A molecule is a combination of two or more atoms held together in a specific shape by attractive forces.

The simplest molecules contain just two atoms For example, a molecule of hydrogen is made up of two hydrogen

atoms A molecule that contains two atoms is classified as a diatomic molecule Figure 1-3 represents a diatomic hydrogen molecule as two spheres connected together

Figure 1-3

A hydrogen molecule can be represented by connecting two spheres together, with each sphere representing one hydrogen atom

Because chemistry deals mostly with the behavior of molecules, this book emphasizes chemistry's molecular

foundation Throughout this book you will see many figures that represent molecules, with each atom represented

by a sphere Although there are more than 100 different types of atoms, only about 20 are encountered frequently

in our world Many of the molecules described in this book are made up of just 10 different types of atoms:

hydrogen, carbon, nitrogen, oxygen, phosphorus, sulfur, fluorine, chlorine, bromine, and iodine Figure 1-4 shows the color scheme that we use to represent these atoms We introduce other types of atoms as the need arises

Although many substances are colored, individual atoms do not have the colors shown in Figure 1-4

Figure 1-4

Color-coded scale models of 10 types of atoms that appear frequently in this book

The name hydrogen refers to both atoms and molecules To minimize confusion, we refer to atomic hydrogen when we mean

hydrogen atoms and molecular hydrogen when we mean hydrogen molecules.

With practice, you can recognize simple molecules, such as carbon dioxide and water, just by looking at their

models Figure 1-5 shows scale models of a few molecules whose names may be familiar to you The structures of larger molecules, however, are too complex to be recognized at a glance Consequently, chemists have created a short-hand language of symbols, formulas, and equations that convey information about atoms and molecules in a simple manner A specific letter or pair of letters designates each type of atom These symbols, in turn, are

combined into formulas that describe the compositions of more complicated chemical substances Formulas can then be used to write chemical equations that describe how molecules change in chemical reactions

Figure 1-5

Models of seven relatively simple molecules

The Elements

A substance that contains only one type of atom is called a chemical element, and each different element contains

its own specific type of atom Each chemical element has a unique name and number , such as hydrogen

, carbon , oxygen , iron , and uranium The name of an element

may refer to its history or to one of its properties The Romans gave copper its name, cuprum, after the island of

Cyprus, where copper was mined as early as 5000 BC Bromine received its name from the Greek word bromos, meaning “stench.” If you ever work with bromine, you will understand the reason for its name Other examples of

evocative names include xenon (from the Greek xenos, “stranger”), rubidium (from the Latin “dark red”), and

neptunium (named for the planet Neptune)

Copper, one of the first elements to be purified by humans, is used for sculpture worldwide.

Each element is represented by a unique one- or two-letter symbol For example, the symbol for hydrogen is H, oxygen's symbol is O, and nitrogen's symbol is N When two or more elements have names that begin with the

same English letter, all but one of the elemental symbols has a second letter The second letter is always lower

case For example, carbon is C, chlorine is Cl, cobalt is Co, and chromium is Cr Chemists understand that the

symbol for an element represents more than one or two letters Instead, a chemist sees the symbol Ni and

immediately thinks of nickel atoms.

The elements listed in Table 1-1 have symbols derived from their names in other languages Most of these

elements were known in ancient times, so their symbols reflect the Latin language that was dominant when they were named

Table 1-1 Elemental Symbols with Non-English Roots

AntimonySb stibium Latin PotassiumK kalium Latin Copper Cu cuprum Latin Silver Ag argentumLatin Gold Au aurum Latin Sodium Na natrium Latin Iron Fe ferrum Latin Tin Sn stannum Latin Lead Pb plumbum Latin Tungsten W wolfram German Mercury Hg hydrargyrumLatin

Chemical Formulas

A chemical compound is a substance that contains more than one element The relative amounts of the elements

in a particular compound do not change: Every molecule of a particular chemical substance contains a

characteristic number of atoms of its constituent elements For example, every water molecule contains two

hydrogen atoms and one oxygen atom To describe this atomic composition, chemists write the chemical formula

for water as

The chemical formula for water shows how formulas are constructed The formula lists the symbols of all

elements found in the compound, in this case H (hydrogen) and O (oxygen) A subscript number after an element's symbol denotes how many atoms of that element are present in the molecule The subscript 2 in the formula for water indicates that each molecule contains two hydrogen atoms No subscript is used when only one atom is

present, as is the case for the oxygen atom in a water molecule Atoms are indivisible, so molecules always

contain whole numbers of atoms Consequently, the subscripts in chemical formulas of molecular substances are always integers We explore chemical formulas in greater detail in Chapter 3

Molecules vary considerably in complexity Molecular oxygen is made up of two oxygen atoms, so its chemical formula is A carbon monoxide molecule contains one atom of carbon and one atom of oxygen, so its

chemical formula is CO Each molecule of methane, the major constituent of natural gas, contains one carbon

atom and four hydrogen atoms, so its formula is You will encounter still more complicated structures, such

as methanol (three different elements, ) as you progress through this book

Section Exercises

c01-0004.What are the elemental symbols for cerium, cesium, copper, calcium, and carbon?

c01-0005.What are the names of the elements represented by the symbols Zr, Ni, Sn, W, Se, Be, and Au?

c01-0006.Molecular pictures of some common molecules are shown here What are their chemical formulas?

The Periodic Table of the Elements

More chemical reactions exist than anyone can imagine Nevertheless, certain patterns of chemical reactivity have been recognized for more than 100 years These patterns remain valid even though new reactions are always being

discovered Each chemical element has characteristic chemical properties Moreover, certain groups of elements

display similar chemical properties In 1869, the Russian chemist Dmitri Mendeleev and the German chemist Julius Lothar Meyer independently discovered how to arrange the chemical elements in a table so that the elements in each

column have similar chemical properties This arrangement, the periodic table, contains all the known chemical

elements

Arrangement

The periodic table lists all the known elements in numerical order, starting with the lightest (hydrogen) and

proceeding to the heaviest (uranium, among naturally occurring elements) The list is broken into seven rows

Each row is placed below the previous row in a way that places elements with similar chemical properties in the

same column of the table Moving across a row of the periodic table, the elements generally increase in mass and change dramatically in their chemical properties Moving down a column, mass also increases, but the elements

have similar chemical properties

We show in Chapter 2 that the periodic table is based on the structure of atoms rather than on their masses Elemental masses correlate closely with atomic structure, however, so ordering by mass is almost the same as ordering by structure There are

only three exceptions among more than 100 elements.

Figure 1-6 shows the periodic table Notice that rows 6 and 7 are quite long, which makes the table rather

cumbersome For convenience, 14 elements in the sixth row and 14 in the seventh row usually are separated from the rest of the table and placed beneath the main portion, as shown in Figure 1-7 and on the inside front cover of the book This is the most common format for the periodic table

Figure 1-6

Periodic table of the elements with all the elements included in their proper rows and columns

Figure 1-7

Metals, Nonmetals, and Metalloids

The elements can be divided into categories: metals, nonmetals, and metalloids Examples of each appear in

Figure 1-7 Except for hydrogen, all the elements in the left and central regions of the periodic table are metals

Metals display several characteristic properties For example, they are good conductors of heat and electricity and

usually appear shiny Metals are malleable, meaning that they can be hammered into thin sheets, and ductile,

meaning that they can be drawn into wires Except for mercury, which is a liquid, all metals are solids at room

temperature

As Figure 1-8 shows, the nonmetals are found in the upper right corner of the periodic table The properties of

nonmetals are highly variable, but most nonmetals are poor conductors of electricity and heat Running diagonally

across the table between the metals and the nonmetals are six elements that are categorized as metalloids (B, Si,

Ge, As, Sb, and Te) These dull-appearing, brittle solids are sometimes called semiconductors because they

conduct electricity better than nonmetals but not as well as metals Silicon and germanium are used in the

manufacture of semiconductor chips in the electronics industry

Figure 1-8

The six metalloids occupy a diagonal region of the periodic table between the metals and the nonmetals

Periodic Properties

The first column of the periodic table, Group 1, contains elements that are soft, shiny solids These alkali metals

include lithium, sodium, potassium, rubidium, and cesium At the other end of the table, fluorine, chlorine,

bromine, iodine, and astatine appear in the next-to-last column These are the halogens, or Group 17 elements

These four elements exist as diatomic molecules, so their formulas have the form A sample of chlorine

appears in Figure 1-7 Each alkali metal combines with any of the halogens in a 1:1 ratio to form a white

crystalline solid The general formula of these compounds is , where A represents the alkali metal and X

represents the halogen ( , LiBr, CsBr, KI, etc.)

The elements in the second column of the table (Group 2) are the alkaline earth metals These resemble the alkali

metals in their appearance, but they have different chemical properties For example, each of these metals

combines with the halogens in a 1:2 ratio ( , , , etc.) Each also reacts with atmospheric oxygen to form a solid with the formula (BaO, CaO, etc.).

The last column of the periodic table contains the noble gases, or Group 18 elements, all of which occur in nature

as gases With a few exceptions, these elements do not undergo chemical reactions

The elements in Groups 3 through 12 are known as transition metals The elements in rows 6 and 7 that are

normally shown below the rest of the table are the inner transition metals, subdivided into lanthanides (row 6) and

actinides (row 7) All other elements are main group elements.

The periodic table is a useful way to organize chemical properties To help you see the patterns, the periodic table

on the inside front cover of this book highlights the various groups of elements As you learn more about chemical structure and behavior, you will discover the principles that account for similarities and differences in the chemical behavior of the elements

Section Exercises

c01-0007.Predict the formulas of the compounds formed in the reactions between (a) calcium and chlorine; (b)

cesium and iodine; (c) barium and oxygen; and (d) magnesium and fluorine

c01-0008.Boron and fluorine form a compound with the formula BF3 Based on this, suggest formulas for

compounds of aluminum with bromine and gallium with chlorine

c01-0009.Classify each of the following elements as alkali metal, alkaline earth metal, halogen, noble gas,

main group, transition metal, lanthanide, or actinide: aluminum, fluorine, cobalt, phosphorus, krypton, europium, thorium, barium, and sodium

Characteristics of Matter

Matter is anything that has mass and occupies space A sample of matter can contain a single substance or any

number of different substances As already described, the building blocks of most substances are molecules, which in turn are composed of atoms It is convenient to classify samples of matter according to the complexity of their

composition, both at the atomic level and at the macroscopic level

The elements are the simplest form of matter An element contains only one type of atom and cannot be decomposed into other chemical components Of the more than 100 known chemical elements, only a few are found in nature in their pure form Figure 1-7 shows three of these: Diamonds are pure carbon, nuggets of pure gold can be found by panning in the right stream bed, and sulfur is found in abundance in its elemental form

When two or more different chemical elements combine, they form a chemical compound Even though they are

made of more than one type of element, pure chemical compounds are uniform in composition; that is, all samples of

a particular chemical compound contain the same proportions of each element For example, ammonia is a chemical compound that contains the elements nitrogen and hydrogen in a 1:3 atomic ratio A sample of pure always contains nothing but ammonia molecules, each one containing three hydrogen atoms bonded to one nitrogen atom

(Figure 1-9)

Figure 1-9

The composition of a pure substance is homogeneous and invariant A sample of pure ammonia contains

nothing but molecules made of nitrogen atoms and hydrogen atoms in a 1:3 ratio

Although pure elements and pure compounds occur often, both in nature and in the laboratory, matter is usually a

mixture of substances A mixture contains two or more chemical substances Unlike pure compounds, mixtures vary

in composition because the proportions of the substances in a mixture can change For example, dissolving sucrose, table sugar, in water forms a mixture that contains water molecules and sucrose molecules A wide range of mixtures can be prepared by varying the relative amounts of sucrose and water

A sample is homogeneous if it always has the same composition, no matter what part of the sample is examined

Pure elements and pure chemical compounds are homogeneous Mixtures can be homogeneous, too; a homogeneous

mixture usually is called a solution As shown in Figure 1-10, the difference between a pure substance and a

homogeneous mixture can be illustrated using hydrogen and chlorine Under the right conditions, these two elements react in a 1:1 atomic ratio to give diatomic molecules of hydrogen chloride Hydrogen chloride gas is a homogeneous

pure substance and always contains equal numbers of hydrogen atoms and chlorine atoms linked in HCl molecules

Under other conditions, molecular hydrogen and molecular chlorine do not react with each other The two gases form

a homogeneous solution whose composition can be changed by adding more of either substance.

Figure 1-10

A sample of hydrogen chloride gas (a) is homogeneous and has constant composition A mixture of hydrogen gas and chlorine gas is homogeneous but can have different compostions (b and c).

When different portions of a mixture have different compositions, the mixture is said to be heterogeneous For

example, quartz is a pure chemical compound made from silicon and oxygen, and gold is a pure element, but the

lump of quartz containing a vein of gold that appears in Figure 1-11a is a heterogeneous mixture because different

parts of the lump have different compositions

Figure 1-11

(a) Quartz containing a vein of gold is a heterogeneous mixture of two different solid substances (b) A glass

of ice water is a heterogeneous mixture of two different phases of a pure substance

Phases of Matter

Matter can also be categorized into three distinct phases: solid, liquid, and gas An object that is solid has a

definite shape and volume that cannot be changed easily Trees, automobiles, ice, and coffee mugs are all in the

solid phase Matter that is liquid has a definite volume but changes shape quite easily A liquid flows to take on

the shape of its container Gasoline, water, and cooking oil are examples of common liquids Solids and liquids are

termed condensed phases because of their well-defined volumes A gas has neither specific shape nor constant

volume A gas expands or contracts as its container expands or contracts Helium balloons are filled with helium gas, and the Earth's atmosphere is made up of gas that flows continually from place to place Molecular pictures that illustrate the three phases of matter appear in Figure 1-12

The gas of the atmosphere is held in place by gravity, not by the walls of a container.

Figure 1-12

Atomic views of the three different phases of matter

Different phases can be mixed together, even when a substance is pure A glass of water containing ice cubes

(Figure 1-11b) is a heterogeneous mixture containing two separate phases of a homogeneous pure substance

Characteristics of Matter

Decide whether each of the following molecular pictures represents a pure substance, a homogeneous

mixture, or a heterogeneous mixture Tell whether the sample is a solid, a liquid, or a gas

Strategy

Apply the characteristics described in the preceding paragraphs Each circle in the figure represents an

atom Different colors distinguish one type of atom from another

Solution

The sample on the left contains a collection of eight diatomic molecules All the molecules have the

same composition, so this is a pure substance The molecules in the sample are distributed evenly

through the entire volume of the container This is the defining characteristic of a gas

In the sample on the right, four atoms of one type are distributed evenly through a larger collection of

atoms of a second type This is a homogeneous mixture Notice that the sample is spread across the

bottom of the container, but it is also confined to a specific volume These features identify the sample as

a liquid

Do the Results Make Sense?

Compare the pictures in the example with those in Figure 1-12 to see that the first sample has the distribution characteristic of a gas, while the second sample has the distribution characteristic of a liquid

Extra Practice Exercise

c01-0146.Draw a molecular picture of dry air, which is a gaseous mixture of 80% molecular nitrogen

and 20% molecular oxygen

Transformations of Matter

When the temperature rises above 100 °C, water changes from a liquid into a gas (steam) Similarly, lava is rock that has been heated sufficiently to convert it to the liquid phase, and “dry ice” is carbon dioxide that has been

cooled enough to change it from the gas phase to the solid phase

Figure 1-13

Water can exist as a solid (ice and snow), as a liquid, and as a gas (steam)

A process that changes the properties of a substance is a transformation Transformations of matter are either

physical or chemical In a physical transformation, physical properties change but the substance's chemical

nature remains the same For example, when water freezes, it undergoes a physical transformation because the

chemical makeup of ice is the same as that of liquid water: Both ice and liquid water contain molecules made up

of two atoms of hydrogen and one atom of oxygen Another physical transformation is the dissolving of a sugar cube in a cup of hot coffee Although the sugar mixes uniformly with the coffee in this process, the chemical

nature of each remains unchanged The Tools for Discovery Box describes some atomic-level transformations

A chemical transformation, on the other hand, produces new substances For example, when magnesium metal

burns in air, elemental magnesium metal and molecular oxygen combine chemically to form magnesium oxide, a white solid containing Mg and O atoms in 1:1 ratio This is a chemical transformation that rearranges the atoms in

Mg and to yield MgO

Tools for Discovery: Atomic-Level Microscopy

Atoms and molecules are much too tiny to see, even with the most powerful light microscopes In recent

years, however, scientists have developed a set of immensely powerful magnifying techniques that make it

possible to visualize how individual atoms are arranged in solids These techniques are scanning tunneling microscopy (STM) and atomic force microscopy (AFM).

Scanning tunneling microscopy and atomic force microscopy generate images of atoms and molecules

using highly miniaturized lever arms with atomically sharp tips—like phonograph needles scaled down to the atomic level These probes respond to the contours of individual atoms In atomic force microscopy, the tip can be moved extremely precisely across a surface The tip responds to individual atoms on that surface, moving up and down by tiny amounts as it passes over each atom A laser beam (like the scanner at a

supermarket checkout stand) focused on the edge of the lever arm detects this tiny motion, and an optical

detector creates an image of the surface

At their highest sensitivities, STM and AFM generate images that show how atoms are arranged on the

surfaces they probe At first, scientists used these tools to explore how atoms are arranged on surfaces The example below shows individual atoms on the surface of nickel metal

Ni metal surface

Increasingly, scientists use STM and AFM to explore chemical reactions For example, molecules often

undergo chemical reactions when they “stick” to a metal surface Scanning tunneling microscopy can be

used to monitor the atomic changes that take place during such chemical reactions It is even possible to use STM to manipulate individual atoms The figures at the bottom of this box show how iron atoms can be

arranged on the surface of copper metal to make an “atomic corral.”

Whereas STM is best suited for imaging atoms, atomic force microscopy is more appropriate for larger

structures The following image shows a strand of DNA The two blue regions of the figure are protein

molecules bound to the DNA

A strand of DNAThe next generation of molecular probes may be able to look within atoms to image their underlying

structure These probes combine magnetic resonance imaging (MRI) with the atomic force microscope

Physicians use MRI to do brain scans to pinpoint tumors or blood clots By coupling MRI with an atomic force scanner and cooling the sample to extremely low temperatures, it may be possible to create images of the interiors of individual atoms

A complete “corral” of iron atoms

Section Exercises

c01-0010.Decide whether each of the following molecular pictures represents a single substance or a mixture

Identify any mixture as homogeneous or heterogeneous If the picture represents a single substance, tell whether it is an element or a chemical compound

c01-0011.Classify each of the following as a pure substance, a solution, or a heterogeneous mixture: a cup of

coffee, a lump of sugar, seasoned salt, a silver coin, and seawater

c01-0012.Classify each of the following as a physical or a chemical transformation: melting snow, burning

coal, chopping wood, and digesting food

Measurements in Chemistry

The knowledge that allows chemists to describe, interpret, and predict the behavior of chemical substances is gained

by making careful experimental measurements The properties of a sample can be divided into physical properties, which can be measured without observing a chemical reaction, and chemical properties, which are displayed only

during a chemical transformation Physical properties include familiar attributes such as size, color, and mass Some chemical properties also are familiar to us As examples, bleach reacts chemically with many colored substances to destroy their colors, and molecular oxygen reacts chemically with many fuels to generate heat

Physical Properties

Length, area, and volume measure the size of an object Length refers to one dimension, area refers to two

dimensions, and volume refers to three dimensions of space Size measurements require standard measuring

devices, such as rulers or measuring cups Figure 1-14 shows some standard laboratory equipment for measuring volume

Figure 1-14

Some common laboratory glassware for measuring volume

In addition to its volume, every object possesses a certain quantity of matter, called its mass Mass

measurements are particularly important in chemistry Consequently, highly accurate mass-measuring machines,

called analytical balances, are essential instruments in chemistry laboratories Analytical balances work by

comparing forces acting on masses A modern balance, whether in a delicatessen or in a chemistry laboratory

(Figure 1-15), compares forces quickly and automatically, providing a digital readout of an object's mass

Figure 1-15

Masses are determined using balances that compare two forces (a) Delicatessen balance (b) Students

wearing protective glasses while using analytical balance

The process of determining mass is called weighing Mass and weight are related, but they are not the same property Mass is a

fundamental characteristic of an object, whereas weight results from gravitational force acting on an object's mass In outer

space, objects have mass but no weight because there is no gravitational force.

Chemists measure time because they want to know how long it takes for chemical transformations to occur Some chemical reactions, such as the conversion of green plants into petroleum, may take millions of years Other chemical processes, such as an explosion of dynamite, are incredibly fast Whereas wristwatches typically measure time only to the nearest second, chemists have developed instruments that make it possible to study processes that occur in less than 0.000 000 000 000 01 second.

Most of us associate temperature with the concepts of hot and cold More accurately, however, temperature

is the property of an object that determines the direction of heat flow Heat flows naturally from a warm object to a cool object, from higher temperature to lower temperature Heat is a form of energy, and because energy changes

in chemical systems have important consequences, chemists are interested in temperature changes that occur

during chemical transformations

All experimental sciences rely on quantitative measurements of properties Every measurement gives a numerical

result that has three aspects: a numerical magnitude; an indicator of scale, called a unit; and a precision Each

aspect is essential, and all three must be reported to make a measurement scientifically valuable

Magnitude

The magnitudes of experimental values in science range from vanishingly small to astronomically large, as

summarized in Figure 1-16 To simplify manipulating very large and very small numbers, we use powers of ten,

also called scientific notation (see Appendix A for a review) For example, the diameter of a carbon atom is 0.000

000 0003 meters (symbol: m) This cumbersome number can be simplified by the use of scientific notation:

Figure 1-16

Dimensions of objects known to humans span more than 30 orders of magnitude, from interstellar distances

to the diameter of a hydrogen nucleus Scientists routinely study objects whose sizes extend far beyond the narrow range encountered in daily life

Physicists, for example, study atomic nuclei measuring m across, and astronomers study our universe,

which spans about m Chemists are most often interested in matter on the smaller side of this range Length measurements in the laboratory vary from meters to subatomic sizes, m To further simplify the use of very large and very small numbers, scientists use prefixes that change the unit sizes by multiples of 10 For instance, the prefix pico means “ ” The symbol for picometer is pm The diameter of a carbon atom is ,

which is 300 pm: The most common magnitude prefixes are listed in Table 1-2.

Table 1-2 Frequently Used Scientific Prefixes for Magnitudes

PrefixSymbolNumber Exponential Notation

tera T 1,000,000,000,000 giga G 1,000,000,000 mega M 1,000,000 kilo k 1000

centi c 0.01 milli m 0.001 micro μ 0.000 001 nano n 0.000 000 001 pico p 0.000 000 000 001

Units

The units associated with a numerical value are just as important as the value itself If a recipe instruction reads,

“Add 1 of sugar,” that instruction is useless The unit of measure, such as “teaspoon,” must also be included It is

essential to include a unit with every experimental value.

A unit of measurement is an agreed-upon standard with which other values are compared Scientists use the meter

as the standard unit of length The meter was originally chosen to be times the length of a line from the

North Pole to the equator Volume can be measured in pints, quarts, and gallons, but the scientific units are the

cubic meter and the liter Temperature can be measured in degrees Fahrenheit (°F), degrees Celsius (°C), or

kelvins (K).

The international scientific community prefers to work exclusively with a single set of units, the Système

International (SI), which expresses each fundamental physical quantity in decimally (power of 10) related units

The seven base units of the SI are listed in Table 1-3 The SI unit for volume is obtained from the base unit for

length: A cube that measures 1 meter on a side has a volume of 1 cubic meter

Table 1-3 Base SI Units

Length Meter m

Temperature Kelvin K

Unit Conversions

It is often necessary to convert measurements from one set of units to another As an everyday example, travelers between the United States and Canada need to be able to convert between miles and kilometers Chemists

frequently need to convert volumes from one unit to another The SI unit of volume is the cubic meter, but

chemists usually work with much smaller volumes Hence chemists often express volume using the liter (L),

which is defined to be exactly Another volume unit in common use is the milliliter (mL), or L The milliliter is the same as the cubic centimeter (cm3) Figure 1-17 illustrates these volume measurements

To convert from one unit to another, we multiply by the ratio that leads to an appropriate cancellation of units For example, a volume of exactly two liters is expressed in quarts as follows:

The correct conversion ratio leads to cancellation of unwanted units

A unit equality may link SI units and non-SI units ( ), decimally related units

, or base units and derived units Some of the more common unit equalities are given on the inside back cover of this text Examples and treat unit conversions

Unit Conversions

Two rock climbers stand at the bottom of a rock face they estimate to be 155 feet high Their rope is 65 m long Is the rope long enough to reach the top of the cliff?

Strategy

The team must compare the English measure for the height of the cliff and the SI measure for the length

of their rope Is 65 m more or less than 155 feet? The equivalence between feet and meters can be found

on the inside back cover of the book:

Solution

The length of the rope is given in meters To set up a ratio that converts meters to feet, we need to have

Now multiply the length of the rope by the conversion ratio: The two climbers have plenty of rope to climb the rock face.

Does the Result Make Sense?

Notice that we multiply the length of the rope in meters by feet/meter so that meters cancel Multiplying the height of the cliff by the same conversion ratio would have given nonsensical units: For a quick estimate, realize that there are a bit over three feet in one meter, so 65 m is a bit more than 195 ft

Extra Practice Exercise

c01-0147.A bottle of lemon juice concentrate contains 24 fluid ounces What is this volume in liters?

(There are 32 fluid ounces in 1 quart.)

Multiple Unit Conversions

The speed limit on many highways in the United States is 55 miles/hr What is this speed limit in SI units?

The speed limit in miles per hour is given To cancel miles, multiply by a ratio that has miles in the

denominator To obtain meters, use a ratio that has meters in the numerator We convert miles to

kilometers and then convert kilometers to meters:

Notice that the conversion ratios are set up to cancel the unwanted units.

This completes the conversion into SI units of length Now convert from hours to seconds, following a similar procedure Because hours appear in the denominator, multiply by a ratio that has hours in the

numerator:

Does the Result Make Sense?

The units are correct, and 25 m/s is a reasonable speed for an automobile

Extra Practice Exercise

c01-0148.A hybrid automobile gets 45 miles per gallon of gasoline in city traffic What is this

mileage in km/L?

The Fahrenheit scale, in which water freezes at 32 °F and boils at 212 °F, is still in common use in the United States, but scientists rarely use the Fahrenheit scale The formula for converting temperature from Fahrenheit to Celsius is

Scientists use two units for temperature, the Celsius (°C) scale and the Kelvin (K) scale These scales are shown

schematically in Figure 1-18 Unlike other scientific units, the unit size of the Celsius and Kelvin scales is the

same, but their zero points differ For both scales, the difference in temperature between the freezing and boiling points of water is defined to be 100 units However, the temperature at which ice melts to liquid water is 0 °C and 273.15 K

Figure 1-18

Celsius and Kelvin temperature scales Kelvins and degrees Celsius have the same unit size but different

zero points “Room temperature” is typically about 295 K (22 °C)

The conversion between kelvins and degrees Celsius is straightforward because their temperature (T) units are the same size A temperature change of 1 °C is the same as a temperature change of 1 K To convert from one scale to the other, add or subtract 273.15:

Precision and Accuracy

The exactness of a measurement is expressed by its precision This concept can be explained with an example

Suppose three swimmers are discussing the temperature of a swimming pool The first dips a finger in the water and says that the temperature is “about 24 °C.” The second examines an immersed pool thermometer and reports the temperature to be 26 °C The third swimmer, who has been monitoring daily variations in the pool's

temperature, uses a portable precision digital thermometer and reports, “According to my precision thermometer, the pool temperature is 25.8 °C.”

The swimmers have measured the water temperature using different measuring devices with different levels of

precision The first swimmer's “finger test” is precise to about 3 °C: (read “twenty-four plus or minus three degrees”) The pool thermometer gives a reading that is precise to the nearest degree:

The precision thermometer measures to the nearest tenth of a degree:

An actual temperature between 25.7 and 25.9 °C falls within the precision ranges of all three measurements, so all three are correct to within their stated limits of precision

Whereas precision describes the exactness of a measurement, accuracy describes how close a measurement is to

the true value Figure 1-19 illustrates the difference between precision and accuracy An expert archer shoots

arrows with high precision, all the arrows hitting the target near the same spot (a and b) Under calm conditions, the expert is also quite accurate, shooting all the arrows into the bull's-eye (a) In a strong wind, however, the

archer may be precise but not accurate, regularly missing the bull's-eye (b) In contrast, a novice shooter is not

very precise (c and d): The arrows scatter over a large area The novice usually is not accurate (c), but an

occasional arrow may accurately find the bull's-eye (d).

Figure 1-19

Patterns of arrows striking a target illustrate the notions of precision and accuracy (a) is precise and

accurate (b) is precise but inaccurate (c) is neither precise nor accurate (d) is an imprecise pattern with

one accurate shot

The goal of any measurement in science is to be as accurate as possible However, determining the accuracy of a measurement is much harder than determining precision An archer can examine the target to see if the arrows

found their mark, but a scientific researcher studying a new phenomenon does not know the correct value A

scientist can assess precision by making additional measurements to find out how closely repeated measurements agree with one another In contrast, assessing accuracy requires careful attention to the design of an experiment and the instruments used in that experiment

Significant Figures

Scientific measurements should always include both magnitude and precision Indicating the precision limits with

a plus/minus (±) statement is cumbersome, particularly when many numerical values are reported at one time

Scientists have agreed to simplify the reporting of precision Experimental measurements are written so that there

is an uncertainty of up to one unit in the last reported digit For example, a temperature reported as 26 °C is greater than 25 °C but less than 27 °C, or A temperature reported as 25.8 °C is The

number of digits expressed in a numerical value is called the number of significant figures The value 26 has two

significant figures, whereas 25.8 has three significant figures

Zeros can present a problem when determining the precision of a numerical value This is because zeros are

needed both to locate the decimal point and to express precision From the statement that the sun is 93,000,000

miles from the Earth, we cannot tell whether the measurement is precise to eight significant figures or whether the zeros are there only to put the decimal point in the right place Scientific notation eliminates this ambiguity

because a power of 10 locates the decimal point, leaving us free to indicate the precision by the number of digits For example, a distance of miles means that the number is precise to miles; the number has two significant figures If this distance is known to a precision of miles, it is written as

miles, with three significant figures Zeros at the end of a number with a decimal point always indicate increased precision Writing “0.010” in scientific notation clarifies that this number has two significant

To determine how many significant figures there are in a particular numerical value, read the number from left to right and count all the digits, starting with the first digit that is not zero To avoid ambiguities about trailing zeros,

we place a decimal point after a value when its trailing zeros are significant For example, “110” has only two

significant figures ( ), whereas “110.” has three significant figures ( ) Here are some

examples of significant figures:

500 1 significant figure 505 3 significant figures

0.051 significant figure 3 significant figures

55 2 significant figures5.000 4 significant figures

50 2 significant figures505.0 4 significant figures

Section Exercises

c01-0013.Convert the following measurements to scientific notation and express in base SI units: 0.000 463

L, 17,935 km, and 260,000 hr (precise to three significant figures)

c01-0014.One light-year is the distance light travels in exactly one year The speed of light is

miles/hr Express the speed of light and the length of one light-year in SI units

c01-0015.Convert each of these measurements to SI units: 155 pounds (mass of a typical person), 120.0

yards (full length of a football field), 39 °C (body temperature of someone with a slight fever), and 365.2422 days (length of 1 year)

Calculations in Chemistry

During an experiment, a chemist may measure physical quantities such as mass, volume, and temperature Usually the chemist seeks information that is related to the measured quantities but must be found by doing calculations In later chapters we develop and use equations that relate measured physical quantities to important chemical properties Calculations are an essential part of all of chemistry; therefore, they play important roles in much of general

chemistry The physical property of density illustrates how to apply an equation to calculations

Density

The density of an object is its mass (m) divided by its volume (V):

The symbol for density is the Greek lower-case letter rho, ρ.

Every pure liquid or solid has a characteristic density that helps distinguish it from other substances To give one example, the density of pure gold is 19.3 g/cm3, whether the sample is a nugget in a miner's pan or an ingot in a bank vault Pyrite, an iron compound that resembles gold, has a much lower density of 5.0 g/cm3 Table 1-4 lists the densities of several common solids and liquids The densities of liquids and solids range from 0.1 to 20 g/cm3 This 200-fold variation is readily apparent to us and dictates how different materials are used in applications in

which density is important Fishermen use floats and sinkers made of low- and high-density materials,

respectively, as Figure 1-20 illustrates Low-density cork is used for floats, because low-density materials float in water High-density lead is used for sinkers, because high-density materials sink.

Table 1-4 Densities of Some Solids and Liquids

Aluminum (Al, solid) 2.70 Acetic acid ( , liquid) 1.05 Bromine ( , liquid)3.10 Acetone ( , liquid) 0.791 Copper (Cu, solid) 8.96 Benzene ( , liquid) 0.885 Gold (Au, solid) 19.3 Chloroform ( , liquid) 1.49 Iodine ( , solid) 4.93 Cork (solid) 0.24 Iron (Fe, solid) 7.87 Diethyl ether ( , liquid)0.714 Lithium (Li, solid) 0.532 Ethanol ( , liquid) 0.785 Lead (Pb, solid) 11.34 Octane ( , liquid) 0.703 Mercury (Hg, liquid) 13.55 Quartz ( , solid) 2.65 Silver (Ag, solid) 10.50 Sodium chloride (NaCl, solid) 2.165 Sodium (Na, solid) 0.968 Water ( , liquid) 1.00 Titanium (Ti, solid) 4.54 Wood (balsa, solid) 0.12 Zinc, (Zn, solid) 7.14 Wood (pine, solid) 0.35–0.50

Figure 1-20

Cork has a lower density than water, so corks of all sizes float on water Lead has a higher density than

water, so lead pieces of all sizes sink in water

Unlike mass and volume, density does not vary with the amount of a substance Notice in Figure 1-20 that all the corks float, regardless of their sizes Notice also that all the pieces of lead sink, regardless of their sizes Dividing a sample into portions changes the mass and volume of each portion but leaves the density unchanged A property

that depends on amount is called extensive Mass and volume are two examples of extensive properties A

property that is independent of amount is called intensive Density and temperature are intensive properties.

Mass and volume often can be measured easily, and density is then calculated using Equation The equation can also be rearranged to find an object's volume or mass, as Example illustrates

Using Density

A diamond is a pure sample of the element carbon The tabulated density of diamond is 3.51 g/cm3

Jewelers use a unit called a carat to describe the mass of a diamond: What is the

volume of the stone in a 2.00-carat diamond engagement ring?

Strategy

Given the diamond's mass and density, we are asked to find its volume Rearranging the density equation makes this possible:

Solution

A list of the information given in the problem allows us to determine what to substitute into the equation:

In the volume equation, density and mass must have consistent units Thus mass must be converted from carats to grams using a conversion ratio that eliminates the unwanted unit:

Now substitute into the equation for volume: