U N I T O N E

The Life of a Cell

Single cells can be complex, independent organisms, such as these twociliates, of the kingdom Protista. A large Euplotes (about 300 mm in length)prepares to eat a much smaller Paramecium. Both are covered with cilia,short, beating structures used to move and ingest prey.

“If there is magic on this planet, it is contained in water.”

Loren Eiseley in The Immense Journey (1957)

Could something this delicious possibly be good for you? Chocolate, made from seeds foundinside cacao pods (inset), contains high levels of antioxidants.

Case Study: Health Food?1 What Are Atoms?

Atoms, the Basic Structural Units of Matter, AreComposed of Still Smaller Particles

2 How Do Atoms Interact to Form Molecules?Atoms Will Interact with Other Atoms Only When ThereAre Vacancies in Their Outermost Electron ShellsCharged Atoms Called Ions Interact to Form Ionic BondsUncharged Atoms Can Become Stable by SharingElectrons, Forming Covalent Bonds

Hydrogen Bonds Are Weaker Electrical AttractionsBetween or Within Molecules with Polar Covalent Bonds

3 Why Is Water So Important to Life?Water Interacts with Many Other MoleculesWater Molecules Tend to Stick TogetherWater Can Form H+ and OH– IonsWater Moderates the Effects of Temperature ChangesWater Forms an Unusual Solid: Ice

Case Study Revisited: Health Food?

2 Atoms, Molecules, and Life

21

A T A G L A N C E

CASESTUDYCASESTUDYCASESTUDYCASESTUDYC

When 35 million people in the UnitedStates gave their loved ones boxes of

chocolates last Valentine’s Day, they knewthey were giving sweet comfort—but healthfood? Sometimes described as “sinfully deli-cious,” chocolate has often been a source ofguilt for those who indulge (or overindulge) init. Chocolate candy is certainly a significantsource of fat and sugar calories, but recentresearch suggests that chocolate itself—the

dark, bitter powder made from the seeds with-in cacao pods (see inset) may also be a signif-icant source of protective molecules. Medicalscientists have known for some time thatmany things that go wrong with our bodiescan be traced to destructive molecules calledfree radicals. Many free radicals contain oxy-gen in a form that reacts strongly with, anddamages, various biological molecules andtheir cellular structures. This process is called

oxidative stress. Oxidative stress tears up cellmembranes, breaks down DNA, and destroysenzymes, resulting in many aspects of aging,cancer, heart disease, and nervous systemdisorders. Unfortunately, oxidative stress is afact of life, because as our cells use energy,they naturally produce free radicals. Sowhat’s a person to do? Well, maybe eatchocolate! n

CASESTUDYHealth Food?

a What Are Atoms?

Atoms, the Basic Structural Units of Matter, Are Composed of Still Smaller Particles

If you took a diamond (a form of carbon) and cut it intopieces, each piece would still be carbon. If you couldmake finer and finer divisions, you would eventuallyproduce a pile of carbon atoms. Atoms are the funda-mental structural units of matter. Atoms themselves,however, are composed of a central atomic nucleus(often called simply the nucleus; plural, nuclei, but don’tconfuse it with the nucleus of a cell). The nucleus con-

tains two types of subatomic particles of equal weight:positively charged protons and uncharged neutrons.Subatomic particles called electrons orbit the atomic nu-cleus (Fig. 2-1). Electrons are lighter, negatively chargedparticles. An atom by itself has an equal number of elec-trons and protons and is therefore electrically neutral.

There are 92 types of atoms that occur naturally. Eachtype of atom forms the structural unit of a different ele-ment. An element is a substance that can neither be bro-ken down nor converted to other substances by ordinarychemical means. The number of protons in the nucleus,called the atomic number, is a characteristic of each ele-ment. For example, every hydrogen atom has one protonin its nucleus, every carbon atom has six protons, and

22 Chapter 2 Atoms, Molecules, and Life

every oxygen atom has eight. Each element has uniquechemical properties based on the number and configura-tion of its subatomic particles. Some elements, such asoxygen and hydrogen, are gases at room temperature;others, such as lead, are extremely dense solids. Most el-ements are quite rare, and relatively few are essential tolife on Earth. Table 2-1 lists the most common elementsin the universe, on Earth, and in the human body. Noticehow differently these elements are distributed.

Atoms of the same element may have different num-bers of neutrons; when this occurs, the atoms are calledisotopes of each other. Some, but not all, isotopes are ra-dioactive; that is, they spontaneously break apart, form-ing different types of atoms and releasing energy in theprocess. Radioactive isotopes are extremely useful as“labels” in studying biological processes (see “ScientificInquiry: Radioactivity in Research”).

Electrons Orbit the Nucleus at Fixed Distances, Forming Electron Shells That Correspond to Different Energy LevelsAs you may know from experimenting with a magnet,like poles repel each other and opposite poles attracteach other. In a similar way, electrons repel one anoth-er, owing to their negative electrical charge, and theyare drawn to the positively charged protons of the nu-cleus. However, because of their mutual repulsion, onlylimited numbers of electrons can occupy the space clos-est to the nucleus. Large atoms can accommodate manyelectrons because their electrons orbit at increasing dis-tances from their nucleus. The electrons orbit through athree-dimensional space; the orbits, which correspond

to different energy levels, are called electron shells (Figs.2-1 and 2-2, p. 24).

The electron shell closest to the atomic nucleus is thesmallest and can hold only two electrons. The secondshell can hold up to eight electrons. The electrons in anatom normally fill the shell closest to the nucleus andthen begin to occupy the next shell. Thus, a carbonatom, with six electrons, has two electrons in the firstshell, closest to the nucleus, and four electrons in its sec-ond shell (see Fig. 2-2). Nuclei and electron shells playcomplementary roles in atoms. Nuclei (assuming theyare not radioactive) provide stability, while the electronshells allow interactions, or bonds, with other atoms.Nuclei resist disturbance by outside forces. Ordinary

Table 2-1 Common Elements Important in Living Organisms

Atomic Percent in Percent in Percent in Element Symbol Numbera Universeb Earthb Human Bodyb

Hydrogen H 1 91 0.14 9.5

Helium He 2 9 Trace Trace

Carbon C 6 0.02 0.03 18.5

Nitrogen N 7 0.04 Trace 3.3

Oxygen O 8 0.06 47 65

Sodium Na 11 Trace 2.8 0.2

Magnesium Mg 12 Trace 2.1 0.1

Phosphorus P 15 Trace 0.07 1

Sulfur S 16 Trace 0.03 0.3

Chlorine Cl 17 Trace 0.01 0.2

Potassium K 19 Trace 2.6 0.4

Calcium Ca 20 Trace 3.6 1.5

Iron Fe 26 Trace 5 Trace

aAtomic number = number of protons in the atomic nucleus.

bApproximate percentage of atoms of this element, by weight, in the universe, in Earth’s crust, and in the human body.

(a)

Hydrogen (H)

electronshell

nucleus

e-

p+

(b)

Helium (He)

nn

e-

p+

p+

e-

Figure 2-1 Atomic modelsStructural representations of the two smallest atoms, (a) hydrogenand (b) helium. In these simplified models, the electrons are repre-sented as miniature planets, circling in specific orbits around a nu-cleus that contains protons and neutrons.

ME

DIA

TU

TO

R

2.1

Inte

racti

ve

Ato

ms

What Are Atoms? 23

Scientific InquiryRadioactivity in Research

As you read this text, you will encounter many statements thatmay cause you to wonder, How do they know that? How dobiologists know that DNA is the genetic material of cells (Chap-ter 9)? How do paleontologists measure the ages of fossils(Chapter 17)? How do botanists know that sugars made inplant leaves during photosynthesis are transported to otherparts of the plant in the sieve tubes of phloem (Chapter 23)?These discoveries, and many more, have been possible onlythrough the use of radioactive isotopes.

Although all atoms of a particular element have the samenumber of protons, the number of neutrons may vary. Neu-trons don’t affect the chemical reactivity of an atom very much,but they add to the atom’s mass, which can be detected by so-phisticated instruments, such as mass spectrometers. Nucleiwith “too many” neutrons break apart spontaneously, ordecay, often emitting radioactive particles in the process. Thoseparticles can also be detected—for example, with Geiger coun-ters. The process by which a radioactive isotope spontaneouslybreaks apart is called radioactive decay.

A particularly fascinating and medically important use ofradioactive isotopes is positron emission tomography, morecommonly known as PET scans (Fig. E2-1a). In one commonapplication of PET scans, a subject is given the sugar glucosethat has been labeled with (that is, attached to) a harmlessradioactive isotope of fluorine. When the nucleus of radioac-tive fluorine decays, it emits two bursts of energy that travelin opposite directions along the same line. Energy detectorsarranged in a ring around the subject look at a “slice” of the

brain, recording the nearly simultaneous arrival of the twoenergy bursts (Fig. E2-1b). A powerful computer then calcu-lates the location within the subject at which the decay musthave occurred and generates a color-coded map of the fre-quency of decays. Because the radioactive fluorine is at-tached to glucose molecules, this map reflects the glucoseconcentrations within the subject’s brain. Since the brainuses this sugar for energy, the more active a brain cell is, themore glucose it uses, and the more radioactivity is concen-trated there. For example, tumor cells (Fig. E2-1c) or brain re-gions in which epileptic seizures originate generally haveexcessively high glucose utilization and show up in PET scansas “hot spots.” Normal brain cells activated by a specificmental task will also have higher glucose demands, whichcan be detected by PET scans.

Biology and medicine have profited immensely from closeinteractions with the other sciences, especially chemistry andphysics. The development of PET scans required close cooper-ation with chemists (developing and synthesizing the radioac-tive probes), physicists (understanding interactions betweenelectrons and short-lived, positively charged particles calledpositrons, as well as the geometry of the resulting energyemissions), and engineers (designing and building the elec-tronic apparatus). Continued teamwork among scientistspromises further advances in both the fundamental under-standing of biological processes and applications in medicineand agriculture.

detector ring

(a) (c)(b)

Subject's head is placed within a ring of detectors.

Radioactive decay releases energetic particles that activate the detectors.

Red indicates the highest radioactivity; blue is least. A malignant brain tumor shows clearly in red.

Figure E2-1 How positron emission tomography works

24 Chapter 2 Atoms, Molecules, and Life

Carbon (C) Oxygen (O) Phosphorus (P) Calcium (Ca)

CaOC P

4e– 6e–

5e–

2e–

8e–

8e–

8e–

2e– 2e–

2e–

2e–

6p+ 8p+ 15p+20p+

6n 8n 16n 20n

Figure 2-2 Electron shells in atomsMost biologically important atoms have at least two shells of electrons. The first shell, closest to the nucleus, can hold twoelectrons; the next shell can hold a maximum of eight electrons. More distant shells can also hold eight electrons each.

sources of energy, such as heat, electricity, and light,hardly affect them at all. Because its nucleus is stable, acarbon atom remains carbon whether it is part of a dia-mond, carbon dioxide, or sugar. Electron shells, howev-er, are dynamic; as you will soon see, atoms bond withone another by gaining, losing, or sharing electrons.

b How Do Atoms Interact to Form Molecules?

Atoms Interact with Other Atoms When There Are Vacancies in Their Outermost Electron Shells

A molecule consists of two or more atoms, of either thesame or of different elements, held together by interac-tions among their outermost electron shells. A sub-stance whose molecules are formed of different types ofatoms is called a compound. Atoms interact with oneanother according to two basic principles:

1. An atom will not react with other atoms when its out-ermost electron shell is completely full or empty. Suchan atom is described as being inert.

2. An atom will react with other atoms when its outer-most electron shell is only partially full. Such atomsare described as reactive.

To demonstrate these principles, consider three atoms:hydrogen, oxygen, and helium (see Fig. 2-1). Hydrogen(the smallest atom) has one proton in its nucleus andone electron in its single (and therefore outermost)electron shell, which can hold up to two electrons. Theoxygen atom has six electrons in its outer shell, whichcan hold eight. In contrast, helium has two protons in itsnucleus, and two electrons fill its single electron shell.Therefore, we predict that hydrogen and oxygen atoms,with partially empty outer shells, should be reactive,while helium atoms, with a full shell, should be stable.We might further predict that hydrogen and oxygenatoms could gain stability by reacting with each other.

Table 2-2 Chemical Bonds

Type of Bond Bond Forms:

Weak Bonds: allow interactions between individual atoms or molecules

Ionic bonds Between positive and negative ions

Hydrogen bonds Between a hydrogen atom involved in a polar covalent bond andanother atom involved in a polar covalent bond

Hydrophobic interactions Because interactions between water molecules exclude hydrophobicmolecules

Strong Bonds: hold atoms together within molecules

Covalent bonds By the sharing of electron pairs; equal sharing produces nonpolarcovalent bonds; unequal sharing produces polar covalent bonds

How Do Atoms Interact to Form Molecules? 25

The single electrons from each of the two hydrogenatoms would fill the outer shell of oxygen, forming H2O:water (see Fig. 2-4c). As predicted, hydrogen can reactreadily with oxygen—in fact, the reaction is an explo-sive one. The space shuttle and many other rockets useliquid hydrogen as fuel to power liftoff. The hydrogenfuel reacts with oxygen, releasing water as a by-product.In contrast, helium, with a full outer shell, is almostcompletely inert. Our case study mentioned free radi-cals. Free radicals are atoms or molecules that lack oneor more electrons in their outer shells, making themhighly reactive.

An atom with an outermost electron shell that is par-tially full can gain stability by losing electrons (empty-ing the shell completely), by gaining electrons (fillingthe shell), or by sharing electrons with another atom(with both atoms behaving as though they had full outershells). The results of losing, gaining, and sharing elec-trons are chemical bonds; attractive forces that holdatoms together in molecules. Each element has chemi-cal bonding properties that arise from the configurationof electrons in its outer shell (Table 2-2). Chemical reac-tions, the making and breaking of chemical bonds toform new substances, are essential for the maintenanceof life and for the working of modern society. Whetherthey occur in a plant cell as it captures solar energy, yourbrain as it forms new memories, or your car’s engine asit guzzles gas, chemical reactions consist of making newchemical bonds and/or breaking existing ones.

Charged Atoms Called Ions Interact to Form Ionic Bonds

Both atoms that have an almost empty outermost elec-tron shell and atoms that have an almost full outermostshell can become stable by losing electrons (emptyingtheir outermost shell) or by gaining electrons (fillingtheir outermost shell).The formation of table salt (sodi-um chloride) demonstrates this principle. Sodium (Na)has only one electron in its outermost electron shell,and chlorine (Cl) has seven electrons in its outer shell—one electron short of being full (Fig. 2-3a). Sodium,therefore, can become stable by losing the electron tochlorine from its outer shell, leaving that shell empty;chlorine can fill its outer shell by gaining the electron.Atoms that have lost or gained electrons, altering thebalance between protons and electrons, are charged.These charged atoms are called ions. To form sodiumchloride, sodium loses an electron and thereby becomesa positively charged sodium ion (Na+); chlorine picksup an electron and becomes a negatively charged chlo-ride ion (Cl–) (Fig. 2-3b,c).

Opposite charges attract; therefore, sodium ions andchloride ions tend to stay near one another. They formcrystals that contain repeating orderly arrangements ofthe two ions (Fig. 2-3c). The electrical attraction be-tween oppositely charged ions that holds them together

in crystals is called an ionic bond. Ionic bonds are weakand easily broken, as occurs when salt is dissolved inwater (see Table 2-2).

Uncharged Atoms Can Become Stable by Sharing Electrons, Forming Covalent Bonds

An atom with a partially full outermost electron shellcan also become stable by sharing electrons with an-other atom, forming a covalent bond. Consider the

Sodium atom (neutral)

N a

17p+11p+11n

11p+11n

Chlorine atom (neutral)

18n

17p+18n

C l

Sodium ion (+) Chloride ion (–)

C l –

C l –

(a)

(b)

(c)

Na+

Na+

__

_

__

_

__ _

_

_

_

_ _

__

__

_

_ _

__

_ _

__

_

_

_ _

__

__

_

__

__

_ _

__

__

_ __

_

_

_

__ _

_

Electron transferred

Attraction betweenopposite charges

An ionic compound: NaCl

Figure 2-3 The formation of ions and ionic bonds(a) Sodium has only one electron in its outer electron shell; chlo-rine has seven. (b) Sodium can become stable by losing an elec-tron, and chlorine can become stable by gaining an electron.Sodium becomes a positively charged ion and chlorine a negative-ly charged ion. (c) Because oppositely charged particles attractone another, the resulting sodium ions (Na+) and chloride ions(Cl–) nestle closely together in a crystal of salt, NaCl.

26 Chapter 2 Atoms, Molecules, and Life

hydrogen atom, which has one electron in a shell builtfor two. A hydrogen atom can become reasonably sta-ble if it shares its single electron with another hydro-gen atom, forming a molecule of hydrogen gas, H2 (Fig.2-4a). Because the two hydrogen atoms are identical,neither nucleus can exert more attraction and capturethe other’s electron. So the two electrons orbit aroundboth nuclei for equal amounts of time, forming a singlecovalent bond; each hydrogen atom behaves almost asif it had two electrons in its shell. Two oxygen atomsalso share electrons equally, producing a molecule ofoxygen gas, O2, with a double covalent bond (Fig.2-4b). In such a bond, each atom contributes two elec-trons. If atoms share three pairs of electrons, a triplecovalent bond is formed, as in nitrogen gas, N2.

All covalent bonds are strong compared with ionicbonds, but some are stronger than others, depending onthe atoms involved (see Table 2-2). Some covalentbonds, such as those in water (H2O; Fig. 2-4c) and carbon

dioxide (CO2), are extremely stable—that is, it takes a lot of energy to break the bonds. Other bonds, such as those in oxygen gas (Fig. 2-4b) or gasoline, are less stable, and come apart more easily. When a chemical reaction occurs in which less stable bonds are brokenand more stable bonds are formed (such as burninggasoline with oxygen to form carbon dioxide and water),energy is released, as we shall describe in Chapter 6.

Most Biological Molecules Utilize Covalent BondingCovalent bonds are crucial to life because the atoms inmost biological molecules are joined by covalentbonds. The molecules in proteins, sugars, bone, and cellulose are formed of atoms held together by cova-lent bonds. Hydrogen, carbon, oxygen, nitrogen, phos-phorus, and sulfur are the most common atoms foundin biological molecules. Except for hydrogen, each ofthese atoms needs at least two electrons to fill its outer-most electron shell and can share electrons with two or

__

(a)

(c)

(b)

Water (H O H or H2O),a polar molecule

Hydrogen (H H or H2),a nonpolar molecule

Oxygen (O O or O2),a nonpolar molecule

+ +__

__ _

__

__ _

____

__

__

__ _ _

__

++

(slightly negative)

(slightly positive)

8p+8n

8p+8n

8p+8n

nonpolar covalentbonding

polar covalentbonding

_ _

Figure 2-4 Covalent bondsElectrons are shared between atoms to form covalent bonds. (a) In hydrogen gas, one electron from each hydrogenatom is shared, forming a single covalent bond. The resulting molecule of hydrogen gas is represented as H–H or H2. (b) In oxygen gas, two oxygen atoms share four electrons, forming a double bond (O=O or O2). (c) Oxygen lacks two electrons to fill its outer shell, so oxygen can make one bond with each of two hydrogen atoms to form water (H–O–H orH2O). Oxygen exerts a greater pull on the electrons than does hydrogen, so the oxygen end of the molecule has a slightnegative charge and the hydrogen end has a slight positive charge. This is an example of polar covalent bonding. Thewater molecule with its slightly charged ends is called a polar molecule.

more other atoms. Hydrogen can form a covalent bondwith one other atom; oxygen and sulfur with two otheratoms; nitrogen with three; and phosphorus and carbonwith up to four. (Phosphorus is unusual; although it hasonly three spaces in its outer shell, it can form up to fivecovalent bonds with up to four other atoms.) This di-versity of bonding arrangements permits biologicalmolecules to be constructed in an almost infinite vari-ety and complexity. Double and triple bonds increasethe variety of shapes and functions of biological mole-cules. Table 2-3 summarizes bonding patterns in biolog-ical molecules.

Polar Covalent Bonds Form When Atoms Share Electrons UnequallyIn hydrogen gas the two nuclei are identical, and theshared electrons spend equal time near each nucleus.Therefore, not only is the molecule as a whole electri-cally neutral, but each end, or pole, of the molecule isalso electrically neutral. Such an electrically symmetri-cal bond is called a nonpolar covalent bond and thecompound formed with such nonpolar bonds is a non-polar molecule, such as hydrogen (H2) or oxygen (O2)(see Fig. 2-4a,b). But electron sharing in covalent bondsis not always equal. In many molecules, one nucleusmay initially have a larger positive charge, and there-fore attract the electrons more strongly, than does theother nucleus. This situation produces a polar covalentbond. Although the molecule as a whole is electricallyneutral, it has charged parts: The atom that attracts theelectrons more strongly then picks up a slightly nega-tive charge (the negative pole of the molecule), and theother atom has a slightly positive charge (the positivepole). In water, for example, oxygen attracts electronsmore strongly than does hydrogen, so the oxygen endof a water molecule is negative and each hydrogen ispositive (see Fig. 2-4c). Water with its charged ends is apolar molecule.

Hydrogen Bonds Are Weaker Electrical Attractions Between or Within Molecules with Polar Covalent Bonds

Because of the polar nature of their covalent bonds,nearby water molecules attract one another. The par-tially negatively charged oxygens of some water mole-cules attract the partially positively charged hydrogensof other water molecules. This electrical attraction iscalled a hydrogen bond (Fig. 2-5; see Table 2-2). As weshall see shortly, hydrogen bonds between moleculesgive water several unusual properties that are essentialto life on Earth.

Hydrogen bonds are common and important in bio-logical molecules, as well as in water. They may occurwhenever polar covalent bonds produce slightly nega-tive and slightly positive charges that then attract oneanother. Both nitrogen and oxygen atoms attract elec-trons more strongly than do hydrogen atoms. Therefore,the nitrogen or oxygen pole of a nitrogen–hydrogen oroxygen–hydrogen bond is slightly negative, and the hy-drogen pole is slightly positive. The resulting polar partsof the molecules can form hydrogen bonds with water,with other biological molecules, or with polar parts ofthe same molecule.Although individual hydrogen bondsare quite weak, many of them working together arequite strong. As we shall see in Chapter 3, hydrogenbonds play crucial roles in shaping the three-dimensionalstructures of proteins. In Chapter 9 you’ll discover theirimportance in DNA .

c Why Is Water So Important to Life?

Water is extraordinarily abundant on Earth, has unusu-al properties, and is so essential to life that it merits spe-cial consideration. Life is very likely to have arisen in

Why Is Water So Important to Life? 27

Table 2-3 Bonding Patterns of Atoms Commonly Found in Biological Molecules

Capacity of Electrons in Number of Covalent Common BondingAtom Outer Electron Shell Outer Shell Bonds Normally Formed Patterns

Hydrogen 2 1 1

Carbon 8 4 4

Nitrogen 8 5 3

Oxygen 8 6 2

Phosphorus 8 5 5

Sulfur 8 6 2

N

O

N

S

C C

N

O

C

P

C

H

28 Chapter 2 Atoms, Molecules, and Life

_

_

hydrogenbonds

O( )

H(+)

H(+)

O( )H

(+)

H(+)

Figure 2-5 Hydrogen bondsThe partial charges on different parts of water molecules produceweak attractive forces called hydrogen bonds (dotted lines) be-tween the hydrogens of one water molecule and the oxygens ofother molecules.

the waters of the primeval Earth. Living organisms stillcontain about 60% to 90% water, and all life dependsintimately on the properties of water. Why is water socrucial to life?

Water Interacts with Many Other Molecules

Water enters into many of the chemical reactions thatoccur in living cells. The oxygen that green plants re-lease into the air is derived from water during photo-synthesis. In manufacturing a protein, fat, nucleic acid,or sugar, your body produces water in the process; con-versely, when you digest proteins, fats, and sugars in thefoods you eat, water is used in the reactions. Why iswater so important in biological chemical reactions?

Water is an extremely good solvent—that is, it is ca-pable of dissolving a wide range of substances, includ-ing protein, salts, and sugars. Water or other solventscontaining dissolved substances are called solutions.Recall that a crystal of table salt is held together by theelectrical attraction between positively charged sodiumions and negatively charged chloride ions (see Fig.2-3c). Because water is a polar molecule, it has positive

and negative poles. If a salt crystal is dropped intowater, the positively charged hydrogen ends of watermolecules will be attracted to and will surround thenegatively charged chloride ions, and the negativelycharged oxygen poles of water molecules will surroundthe positively charged sodium ions. As water moleculesenclose the sodium and chloride ions and shield themfrom interacting with each other, the ions separatefrom the crystal and drift away in the water—and thesalt dissolves (Fig. 2-6).

Water also dissolves molecules held together bypolar covalent bonds. Its positive and negative polesare attracted to oppositely charged regions of dissolv-ing molecules. Ions and polar molecules are termed hy-drophilic (Greek for “water-loving”) because of theirelectrical attraction for water molecules. Many biolog-ical molecules, including sugars and amino acids, arehydrophilic and dissolve readily in water (Fig. 2-7).Water also dissolves gases such as oxygen and carbondioxide. The fish swimming below the ice on a frozenlake rely on oxygen that dissolved before the iceformed, and they release CO2 into solution in thewater. By dissolving such a wide variety of molecules,the watery substance inside a cell provides a suitableenvironment for the countless chemical reactions es-sential to life on Earth.

Molecules that are uncharged and nonpolar, such asfats and oils, usually do not dissolve in water and henceare called hydrophobic (“water-fearing”). Nevertheless,water has an important effect on such molecules. Oils,for example, form globules when spilled into water. Oil

Cl–

Cl–

Na+

Na+

Na+

H

H

H

H

O

O–

Figure 2-6 Water as a solventThe polarity of water molecules allows water to dissolve polar andcharged substances. When a salt crystal is dropped into water,the water surrounds the sodium and chloride ions with oppositelycharged poles of the water molecules. Thus insulated from the at-tractiveness of other molecules of salt, the ions disperse, and thewhole crystal gradually dissolves.

ME

DIA

TU

TO

R

2.2

Wate

r a

nd

Lif

e

Why Is Water So Important to Life? 29

hydrogen bond

hydroxyl group

glucose

water

Figure 2-7 Water dissolves many biological moleculesMany biological molecules dissolve in water because they havepolar parts—for example, OH– (hydroxyl) groups—that can formhydrogen bonds with water molecules. As shown, hydrogen bondscan form between the hydroxyl groups on a glucose molecule (asimple sugar) and surrounding water molecules.

molecules in water disrupt the hydrogen bondingamong adjacent water molecules. When oil moleculesencounter one another in water, their nonpolar surfacesnestle closely together, surrounded by water moleculesthat form hydrogen bonds with one another but notwith the oil. To separate again, the oil molecules wouldhave to break apart the hydrogen bonds that link sur-

rounding water molecules. Thus, the oil molecules re-main together, forming a glistening droplet that floatson the water’s surface. The tendency of oil molecules toclump together is described as a hydrophobic interaction(see Table 2-2). As we shall discuss in Chapter 4, themembranes of living cells owe much of their structureto hydrophobic interactions.

Water Molecules Tend to Stick Together

In addition to interacting with other molecules, watermolecules interact with each other. Because hydrogenbonds interconnect individual water molecules, liquidwater has high cohesion—that is, water molecules havea tendency to stick together. Cohesion among watermolecules at the water’s surface produces surface ten-sion, the tendency for the water surface to resist beingbroken. If you’ve ever experienced the slap and sting ofa belly flop into a swimming pool, you’ve discoveredfirsthand the power of surface tension. Surface tensioncan support fallen leaves, some spiders and water in-sects, and even a running lizard (Fig. 2-8a).

A more important role of cohesion in water occursin the life of land plants. Since a plant absorbs waterthrough its roots, how does the water reach the above-ground parts, especially if the plant is a 100-meter-tallredwood (Fig. 2-8b)? As we shall see in Chapter 23,water molecules are pulled up by the leaves. Waterfills tiny tubes that connect the leaves, stem, and roots.Water molecules that evaporate from the leaves pull

(b)(a)

Figure 2-8 Cohesion among water molecules(a) With webbed feet bearing specialized scales, the basilisk lizard of South America makes use of surface tension,caused by cohesion, to support its weight as it races across the surface of a pond. (b) In giant redwoods, cohesionholds water molecules together in continuous strands from the roots to the topmost leaves even 300 feet (about 100 meters) above the ground.

30 Chapter 2 Atoms, Molecules, and Life

water up the tubes, much like a chain being pulled upfrom the top. The system works because the hydrogenbonds interconnecting water molecules are strongerthan the weight of the water in the tubes, even 100meters’ worth; thus, the water “chain” doesn’t break.Without the cohesion of water, there would be noland plants as we know them, and terrestrial lifewould undoubtedly have evolved quite differently.You may have realized by now that the “commonbond” producing the sting of a belly flop, the ability ofa lizard to run on water, and the movement of waterup a tree is actually the hydrogen bond between watermolecules.

Water exhibits another property, adhesion, a wordthat describes its tendency to stick to polar surfaceshaving slight charges that attract polar water mole-cules. Adhesion helps water move within small spaces,such as the thin tubes in plants that carry water fromroots to leaves. If you stick the end of a narrow glasstube into water, the water will move a short distance upthe tube. Put some water in a narrow glass bud vase ortest tube and you’ll see that the upper surface is curved;water pulls itself up the sides of the glass by its adhe-sion to the surface of the glass and by the cohesionamong water molecules.

Water Can Form H+ and OH– Ions

Although water is generally regarded as a stable com-pound, individual water molecules constantly gain, lose,and swap hydrogen atoms. As a result, at any given timeabout two of every billion water molecules are ion-ized—that is, broken apart into hydrogen ions (H+) andhydroxide ions (OH–):

A hydroxide ion has gained an electron from the hydro-gen ion, giving it a negative charge, while the hydrogenion, which has lost its electron, now has a positivecharge. Pure water contains equal concentrations of hy-drogen ions and hydroxide ions.

In many solutions, however, the concentrations ofH+ and OH– are not the same. If the concentration ofH+ exceeds the concentration of OH–, the solution isacidic. An acid is a substance that releases hydrogenions when it is dissolved in water. When hydrochloricacid (HCl), for example, is added to pure water, almostall of the HCl molecules separate into H+ and Cl–.Therefore, the concentration of H+ greatly exceeds theconcentration of OH–, and the resulting solution is

O O

hydrogen ionhydroxide ionwater(H2O) (OH–) (H+)

+(+)

( – )

H H H

H

acidic. (Many acidic substances, such as lemon juiceand vinegar, have a sour taste. This is because the sour-taste receptors on your tongue are specialized to re-spond to the excess of H+.)

If the concentration of OH+ is greater, the solutionis basic. A base is a substance that combines with hy-drogen ions, reducing their number. If, for instance,sodium hydroxide (NaOH) is added to water, theNaOH molecules separate into Na+ and OH–. TheOH– combine with H+, reducing their number. The so-lution is then basic.

The degree of acidity is expressed on the pH scale(Fig. 2-9), in which neutrality (equal numbers of H+ andOH–) is assigned the number 7. Acids have a pH below7; bases have a pH above 7. Pure water, with equal con-centrations of H+ and OH–, has a pH of 7. Each unit onthe pH scale represents a tenfold change in the concen-tration of H+. Thus, a cola drink (pH = 3) has a concen-tration of H+ 10,000 times that of water (pH = 7)—nowonder it is bad for your teeth!

A Buffer Helps Maintain a Solution at a Relatively Constant pHIn most mammals, including humans, both the cell inte-rior (cytoplasm) and the fluids that bathe the cells arenearly neutral (pH about 7.3 to 7.4). Small increases ordecreases in pH may cause drastic changes in both thestructure and function of biological molecules, leadingto the death of cells or entire organisms. Nevertheless,living cells seethe with chemical reactions that take upor give off H+. How, then, does the pH remain constantoverall? The answer lies in the many buffers found inliving organisms. A buffer is a compound that tends tomaintain a solution at a constant pH by accepting or re-leasing H+ in response to small changes in H+ concen-tration. If the H+ concentration rises, buffers combinewith them; if the H+ concentration falls, buffers releaseH+. The result is that the concentration of H+ is re-stored to its original level. Common buffers in living or-ganisms include bicarbonate (HCO3

–) and phosphate(H2PO4

– and HPO42–), both of which can accept or re-

lease H+, depending on the circumstances. If the bloodbecomes too acidic, for example, bicarbonate acceptsH+ to form carbonic acid:

→(bicarbonate) (hydrogen ion) (carbonic acid)

If the blood becomes too basic, carbonic acid liberateshydrogen ions, which combine with the excess hydrox-ide ions, forming water:

→(carbonic acid) (hydroxide ion) (bicarbonate) (water)

In either case, the result is that the blood pH remainsnear its normal value.

H2O1HCO32OH21H2CO3

H2CO3H11HCO32

Why Is Water So Important to Life? 31

+ _

1-molar hydrochloricacid (HCI)

stomach acidlime juice

lemon juice

"acid rain" (2.5–5.5)vinegar, cola, orange juice,tomatoes

beer

black coffee, tea

normal rain (5.6)urine (5.7)

pure water (7.0)salivablood, sweat (7.4)

seawater (7.8–8.3)

baking soda

phosphate detergentschlorine bleachmilk of magnesia

household ammoniasome detergents(without phosphates)

washing soda

oven cleaner

1-molar sodiumhydroxide (NaOH)

0

1

2

3

4

5

6

7

8

9

10

11

12

13

14

pHvalue

H concentration(moles/liter)

10

10

10

10

10

10

10

10

0

–1

–2

–3

–4

–5

–6

–7

–8

–9

–10

–11

–12

–13

–14

10

10

10

10

10

10

10

+

neutral(H = OH )

incr

easi

ngly

aci

dic

(H

+ >

OH

– )in

crea

sing

ly b

asic

(H+ <

OH

– )

Figure 2-9 The pH scaleThe pH scale expresses the concentration of hydrogen ions in asolution on a scale of 0 (very acidic) to 14 (very basic). Each unit ofchange in pH on the pH scale represents a tenfold change in theconcentration of hydrogen ions. Lemon juice, for example, is about10 times more acidic than orange juice, and the most severe acidrains in the northeastern United States are almost 1000 times moreacidic than normal rainfall. Except for the inside of your stomach,nearly all the fluids in your body are finely adjusted to a pH of 7.4.The color coding corresponds to a common pH indicator dye,bromthymol blue, widely used by aquarium owners to monitor thepH of water for their fish.

Water Moderates the Effects of Temperature Changes

Your body and the bodies of other organisms can sur-vive only within a limited temperature range. As weshall see in Chapter 6, high temperatures may damageenzymes that guide the chemical reactions essential tolife. Low temperatures are also dangerous, because en-zyme action slows as the temperature drops. Subfreez-ing temperatures within the body are usually lethal,because spearlike ice crystals can rupture cells.

Water has important properties that moderate theeffects of temperature changes. These properties helpkeep the bodies of organisms within tolerable tempera-ture limits. Also, large lakes and the oceans have amoderating effect on the climate of nearby land, mak-ing it warmer in winter and cooler in summer. First,some background: Temperature reflects the speed ofmolecules; the higher the temperature, the greater theiraverage speed. Generally speaking, if heat energy en-ters a system, the molecules of that system move morerapidly, and the temperature of the system rises. Indi-vidual water molecules, however, are weakly linked toone another by hydrogen bonds (see Fig. 2-5). Whenheat enters a watery system such as a lake or a livingcell, much of the heat energy goes into breaking hydro-gen bonds rather than speeding up individual mole-cules. Thus, 1 calorie of energy will heat 1 gram of water1 °C, whereas it takes only 0.6 calorie per gram to heatalcohol 1 °C, 0.2 calorie for table salt, and 0.02 caloriefor common rocks such as granite or marble. So the en-ergy required to heat a pound (a pint) of water only 1 °C would raise a pound of rock by 50 °C. If a lizardwants to warm up, it will seek out a rock, rather than apuddle. Because the human body is mostly water, asunbather can absorb a lot of heat energy withoutsending his or her body temperature soaring—andmany hot sunbathers can jump into a swimming pool tocool off without raising the temperature of the watervery much. (The energy required to heat a gram of asubstance by 1 °C is called its specific heat—water has ahigh specific heat.)

Second, water moderates the effects of high temper-atures because it takes a great deal of heat, 539 calo-ries per gram, to convert liquid water to water vapor.This, too, is due to the hydrogen bonds that intercon-nect individual water molecules. For a water moleculeto evaporate, it must move quickly enough to break allthe hydrogen bonds that hold it to the other watermolecules in the solution. Only the fastest-movingwater molecules, carrying the most energy, can breaktheir hydrogen bonds and escape into the air as watervapor. The remaining liquid is cooled by the loss ofthese high-energy molecules. As children rompthrough a sprinkler on a hot summer day, water coatstheir bodies. Heat energy is transferred from their skinto the water and from the water to the vapor as the

REVISITED CASESTUDYREVISITEDCAS

32 Chapter 2 Atoms, Molecules, and Life

water evaporates. Evaporating just 1 gram of watercools 539 grams of a person’s body 1 °C, so water is avery effective coolant. This also means that evaporat-ing perspiration produces a great loss of heat withoutmuch loss of water. (The heat required to vaporizewater is called its heat of vaporization—water’s heat ofvaporization is one of the highest known!)

Third, water moderates the effects of low tempera-tures because an unusually large amount of energy mustbe removed from molecules of liquid water before theyform the precise crystal arrangement of ice. As a result,water freezes more slowly than do many other liquids ata given temperature and loses more heat to the environ-ment in the process. (This property of water is called itsheat of fusion, which is very high.)

Water Forms an Unusual Solid: Ice

Water, of course, will become a solid after prolonged ex-posure to temperatures below its freezing point. Buteven solid water is unusual. Most liquids become denserwhen they solidify, and the solid sinks. Ice is ratherunique because it is less dense than liquid water. Whena pond or lake starts to freeze in winter, the ice stays ontop, forming an insulating layer that delays the freezingof the rest of the water. This insulation allows fish andother lake residents to survive in the liquid water below.If ice were to sink, ponds and lakes in much of NorthAmerica would freeze solid, from the bottom up, duringthe winter, killing fish and underwater plants and mak-ing drinking water far less available to animals.

CASESTUDYREVISITEDHealth Food?

You now know that a free radical is anatom or molecule containing an unfilledouter electron shell, which reacts vigor-ously with other molecules to fill its outershell. Normal cellular activities produce avariety of molecules that contain an oxy-gen atom with an unfilled outer shell.Such molecules react vigorously withother molecules, damaging them in theprocess. If you’ve seen iron turn to rust(iron oxide), you’ve witnessed oxygen’sreactive power. Oxidative stress refers toa kind of ”biological rusting” in whichfree radicals that contain oxygen damagecells. Substances that react with theseoxygen-containing free radicals and ren-der them harmless are called antioxi-

dants. Although cells contain some oftheir own antioxidants, your eating habitscan also make a difference. Fruits andvegetables, particularly those with yellow,orange, and red colors, are good sourcesof antioxidant compounds; vitamins Cand E are antioxidants also. Now, amaz-ingly, researchers have given us an excuseto eat chocolate and feel good about it:cocoa powder contains high concentra-tions of flavenoids, which are powerfulantioxidants. This research is in its earlystages, and no studies have been done todetermine whether high consumption ofchocolate actually reduces the risk ofcancer or heart disease, but there will cer-tainly be no shortage of volunteers for

this research. Although becoming fat byeating too much chocolate candy couldcounteract any positive effects of thecocoa powder itself, slim “chocoholics”have reason to relax and enjoy.

As you will learn in Chapter 8, oxygen isimportant for harvesting the maximumamount of energy from molecules such assugar. Ross Hardison, a researcher atPennsylvania State University, stated elo-quently: “Keeping oxygen under controlwhile using it in energy production hasbeen one of the great compromisesstruck in the evolution of life on Earth.”What did he mean by this?

Summary of Key Concepts

1 What Are Atoms?An element is a substance that can neither be brokendown nor converted to different substances by ordinarychemical means. The smallest possible particle of an ele-ment is the atom, which is itself composed of a central nu-cleus, containing protons and neutrons, and electronsoutside the nucleus. All atoms of a given element have thesame number of protons, which is different from the num-ber of protons in the atoms of every other element. Elec-trons orbit the nucleus in electron shells, at specificdistances from the nucleus. Each shell can contain a fixedmaximum number of electrons. The chemical reactivity ofan atom depends on the number of electrons in its outer-most electron shell: An atom is most stable, and thereforeleast reactive, when its outermost shell is either complete-ly full or empty.

2 How Do Atoms Interact to Form Molecules?Atoms may combine to form molecules.The forces holdingatoms together in molecules are called chemical bonds.Atoms that have lost or gained electrons are negatively orpositively charged particles called ions. Ionic bonds areelectrical attractions between charged ions, holding themtogether in crystals. When two atoms share electrons, cova-lent bonds form. In a nonpolar covalent bond, the twoatoms share electrons equally. In a polar covalent bond,one atom may attract the electron more strongly than theother atom does; in this case, the strongly attracting atombears a slightly negative charge, and the weakly attractingatom bears a slightly positive charge. Some polar covalentbonds give rise to hydrogen bonding, the attraction be-tween charged regions of individual polar molecules or dis-tant parts of a large polar molecule.

3 Why Is Water So Important to Life?Properties of the water molecule that are important to liv-ing organisms include its ability to interact with manyother molecules and to dissolve many polar and chargedsubstances; to force nonpolar substances, such as fat, to as-

sume certain types of physical organization; to participatein chemical reactions; to cohere to itself by using hydrogenbonds between water molecules; and to maintain a fairlystable temperature in the face of wide temperature fluctu-ations in the environment.

Thinking Through the Concepts 33

Key Termsacid p. 30acidic p. 30atom p. 21atomic nucleus p. 21atomic number p. 21base p. 30basic p. 30buffer p. 30calorie p. 31

chemical bond p. 25chemical reaction p. 25cohesion p. 29compound p. 24covalent bond p. 25electron p. 21electron shell p. 22element p. 21hydrogen bond p. 27

hydrophilic p. 28hydrophobic p. 28hydrophobic interaction p. 29ion p. 25ionic bond p. 25isotope p. 22molecule p. 24neutron p. 21nonpolar covalent bond p. 27

pH scale p. 30polar covalent bond p. 27proton p. 21radioactive p. 22solvent p. 28surface tension p. 29

Thinking Through the ConceptsMultiple Choice

1. What is the purest form of matter that cannot be separatedinto different substances by chemical means?a. compoundsb. moleculesc. atomsd. elementse. electrons

2. Which phrase best describes chemical bonds?a. physical bridgesb. attractive forcesc. shared protonsd. atomic reactionse. all of these phrases are equally descriptive

3. When an atom ionizes, what happens?a. It shares one or more electrons with another atom.b. It emits energy as it loses extra neutrons.c. It gives up or takes up one or more electrons.d. It shares a hydrogen atom with another atom.e. none of the above

4. If electrons in water molecules were equally attracted to hy-drogen nuclei and oxygen nuclei, water molecules would bea. more polarb. less polarc. unchangedd. triple bondede. unable to form

5. A covalent bond formsa. when two ions are attracted to one anotherb. between adjacent water molecules, producing surface

tensionc. when one atom gives up its electron to another atomd. when two atoms share electronse. between water molecules and fat globules

6. What is the defining characteristic of an acid?a. It donates hydrogen ions.b. It accepts hydrogen ions.c. It will donate or accept hydrogen ions, depending on

the pH.d. It has an excess of hydroxide ions.e. It has a pH greater than 7.

? Review Questions1. What are the six most abundant elements that occur in liv-

ing organisms?

2. Distinguish among atoms and molecules; elements andcompounds; and protons, neutrons, and electrons.

3. Compare and contrast covalent bonds and ionic bonds.

4. Why can water absorb a great amount of heat with littleincrease in its temperature?

5. Describe how water dissolves a salt. How does this phe-nomenon compare with the effect of water on a hy-drophobic substance such as corn oil?

6. Define acid, base, and buffer. How do buffers reducechanges in pH when hydrogen ions or hydroxide ions areadded to a solution? Why is this phenomenon importantin organisms?

Raloff, J.“Chocolate Hearts.” Science News, March 18, 2000. Describesrecent research indicating that chocolate is high in antioxidants.

Storey, K. B., and Storey, J. M.“Frozen and Alive.” Scientific American,December 1990. By triggering ice formation here, suppressing itthere, and loading up their cells with antifreeze molecules, some an-imals, including certain lizards and frogs, can survive with 60% oftheir body water frozen solid.

34 Chapter 2 Atoms, Molecules, and Life

For More InformationAtkins, P. W. Molecules. New York: Scientific American Library, 1987.

A layperson’s introduction to atoms and molecules, with superb il-lustrations.

Glasheen, J. W., and McMahon, T. A. “Running on Water.” ScientificAmerican, September 1997. Answers the question, “How does thebasilisk lizard run on water?”

Morrison, P., and Morrison, P. Powers of Ten. New York: W. H. Free-man, 1982.A fascinating journey from the universe to the nucleus ofan atom.

Applying the Concepts1. Many “over-the-counter” substances are sold to bring re-

lief from “acid stomach” or “heartburn.” What is thechemical basis for these compounds? Why do they work?

2. Fats and oils do not dissolve in water; polar and ionic mol-ecules dissolve easily in water. Detergents and soaps helpclean by dispersing fats and oils in water so that they canbe rinsed away. From your knowledge of the structure ofwater and the hydrophobic nature of fats, what generalchemical structures (for example, polar or nonpolar parts)must a soap or detergent have, and why?

3. What would the effects be for aquatic life if the density ofice were greater than that of liquid water? What would bethe impact on terrestrial organisms?

4. How does sweating help you regulate your body tempera-ture? Why do you feel hotter and more uncomfortable ona hot, humid day than on a hot, dry day?

Answers to Multiple-Choice Questions1.d2.b3.c4.b5.d6.a

M E D I A T U T O R

Atoms, Molecules, and Life

CD Activities•Activity 2.1: Interactive AtomsEstimated time: 5 minutes

This tutorial provides an introduction to atomicstructure and chemical bonding. An understandingof simple chemistry is critical to understanding biol-ogy. This tutorial will provide the background tolearn about more complex biological molecules.The bonds that hold together large, complex, bio-logically important molecules like DNA are thesame as the bonds we will learn about here.

Activity 2.2: Water and LifeEstimated time: 5 minutes

Water is essential for life. In fact, water may be theonly absolute essential for living systems. This tuto-rial explores the properties of water and how theyrelate to living systems.

Start the MediaTutor Student CD-ROM and enterthe activity number in the Quick Search box to betaken directly to that activity.

Web Investigations•Case Study: Health Food?Estimated time: 10 minutes

Ah, creamy, sweet, delicious chocolate! Most peo-ple love it. Some even think it’s addictive. This exer-cise takes a closer, scientific look at an old favorite.

Go to http://www.prenhall.com/audesirk6, the Aude-sirk Companion Web site. Select Chapter 2 and theWeb Investigation to begin.

MediaTutor 35