the nature of molecules - nova scotia department of...

19

2The Nature of Molecules

Concept Outline

2.1 Atoms are nature’s building material.

Atoms. All substances are composed of tiny particles calledatoms, each a positively charged nucleus around which orbitnegative electrons.Electrons Determine the Chemical Behavior of Atoms.Electrons orbit the nucleus of an atom; the closer anelectron’s orbit to the nucleus, the lower its energy level.

2.2 The atoms of living things are among the smallest.

Kinds of Atoms. Of the 92 naturally occurring elements,only 11 occur in organisms in significant amounts.

2.3 Chemical bonds hold molecules together.

Ionic Bonds Form Crystals. Atoms are linked togetherinto molecules, joined by chemical bonds that result fromforces like the attraction of opposite charges or the sharing ofelectrons.Covalent Bonds Build Stable Molecules. Chemicalbonds formed by the sharing of electrons can be very strong,and require much energy to break.

2.4 Water is the cradle of life.

Chemistry of Water. Water forms weak chemicalassociations that are responsible for much of the organizationof living chemistry.Water Atoms Act Like Tiny Magnets. Because electronsare shared unequally by the hydrogen and oxygen atoms ofwater, a partial charge separation occurs. Each water atomacquires a positive and negative pole and is said to be “polar.”Water Clings to Polar Molecules. Because the oppositepartial charges of polar molecules attract one another, watertends to cling to itself and other polar molecules and toexclude nonpolar molecules.Water Ionizes. Because its covalent bonds occasionallybreak, water contains a low concentration of hydrogen (H+)and hydroxide (OH–) ions, the fragments of broken watermolecules.

A bout 10 to 20 billion years ago, an enormous explo-sion likely marked the beginning of the universe.

With this explosion began the process of evolution, whicheventually led to the origin and diversification of life onearth. When viewed from the perspective of 20 billionyears, life within our solar system is a recent development,but to understand the origin of life, we need to considerevents that took place much earlier. The same processesthat led to the evolution of life were responsible for theevolution of molecules (figure 2.1). Thus, our study of lifeon earth begins with physics and chemistry. As chemicalmachines ourselves, we must understand chemistry tobegin to understand our origins.

FIGURE 2.1Cells are made of molecules. Specific, often simple, combina-tions of atoms yield an astonishing diversity of molecules withinthe cell, each with unique functional characteristics.

weight will be greater on the earth because the earth’s grav-itational force is greater than the moon’s. The atomicmass of an atom is equal to the sum of the masses of itsprotons and neutrons. Atoms that occur naturally on earthcontain from 1 to 92 protons and up to 146 neutrons.

The mass of atoms and subatomic particles is measuredin units called daltons. To give you an idea of just how smallthese units are, note that it takes 602 million million billion(6.02 × 1023) daltons to make 1 gram! A proton weighs ap-proximately 1 dalton (actually 1.009 daltons), as does a neu-tron (1.007 daltons). In contrast, electrons weigh only 1

1840 ofa dalton, so their contribution to the overall mass of an atomis negligible.

20 Part I The Origin of Living Things

AtomsAny substance in the universe that hasmass (see below) and occupies space isdefined as matter. All matter is com-posed of extremely small particlescalled atoms. Because of their size,atoms are difficult to study. Not untilearly in this century did scientistscarry out the first experiments sug-gesting what an atom is like.

The Structure of Atoms

Objects as small as atoms can be“seen” only indirectly, by using verycomplex technology such as tunnelingmicrocopy. We now know a greatdeal about the complexities of atomicstructure, but the simple view putforth in 1913 by the Danish physicistNiels Bohr provides a good startingpoint. Bohr proposed that every atompossesses an orbiting cloud of tinysubatomic particles called electronswhizzing around a core like the plan-ets of a miniature solar system. At thecenter of each atom is a small, verydense nucleus formed of two otherkinds of subatomic particles, protonsand neutrons (figure 2.2).

Within the nucleus, the cluster ofprotons and neutrons is held togetherby a force that works only over shortsubatomic distances. Each proton car-ries a positive (+) charge, and eachelectron carries a negative (–) charge.Typically an atom has one electronfor each proton. The number of protons (the atom’satomic number) determines the chemical character of theatom, because it dictates the number of electrons orbitingthe nucleus which are available for chemical activity. Neu-trons, as their name implies, possess no charge.

Atomic Mass

The terms mass and weight are often used interchangeably,but they have slightly different meanings. Mass refers to theamount of a substance, while weight refers to the forcegravity exerts on a substance. Hence, an object has thesame mass whether it is on the earth or the moon, but its


FIGURE 2.2Basic structure of atoms. All atoms have a nucleus consisting of protons and neutrons,except hydrogen, the smallest atom, which has only one proton and no neutrons in itsnucleus. Oxygen, for example, has eight protons and eight neutrons in its nucleus. Electronsspin around the nucleus a far distance away from the nucleus.

Proton

(Positive charge) (No charge) (Negative charge)

Neutron Electron

Hydrogen1 Proton1 Electron

Oxygen8 Protons8 Neutrons8 Electrons

Isotopes

Atoms with the same atomic number (that is, the same num-ber of protons) have the same chemical properties and aresaid to belong to the same element. Formally speaking, anelement is any substance that cannot be broken down to anyother substance by ordinary chemical means. However, whileall atoms of an element have the same number of protons,they may not all have the same number of neutrons. Atoms ofan element that possess different numbers of neutrons arecalled isotopes of that element. Most elements in nature existas mixtures of different isotopes. Carbon (C), for example,has three isotopes, all containing six protons (figure 2.3).Over 99% of the carbon found in nature exists as an isotopewith six neutrons. Because its total mass is 12 daltons (6 fromprotons plus 6 from neutrons), this isotope is referred to ascarbon-12, and symbolized 12C. Most of the rest of the natu-rally occurring carbon is carbon-13, an isotope with sevenneutrons. The rarest carbon isotope is carbon-14, with eightneutrons. Unlike the other two isotopes, carbon-14 is unsta-ble: its nucleus tends to break up into elements with loweratomic numbers. This nuclear breakup, which emits a signifi-cant amount of energy, is called radioactive decay, and iso-topes that decay in this fashion are radioactive isotopes.

Some radioactive isotopes are more unstable than othersand therefore decay more readily. For any given isotope,however, the rate of decay is constant. This rate is usuallyexpressed as the half-life, the time it takes for one half of theatoms in a sample to decay. Carbon-14, for example, has ahalf-life of about 5600 years. A sample of carbon containing1 gram of carbon-14 today would contain 0.5 gram of car-bon-14 after 5600 years, 0.25 gram 11,200 years from now,0.125 gram 16,800 years from now, and so on. By determin-ing the ratios of the different isotopes of carbon and otherelements in biological samples and in rocks, scientists areable to accurately determine when these materials formed.

While there are many useful applications of radioactivity,there are also harmful side effects that must be considered inany planned use of radioactive substances. Radioactive sub-stances emit energetic subatomic particles that have the po-

tential to severely damage living cells, producing mutations intheir genes, and, at high doses, cell death. Consequently, ex-posure to radiation is now very carefully controlled and regu-lated. Scientists who work with radioactivity (basic re-searchers as well as applied scientists such as X-raytechnologists) wear radiation-sensitive badges to monitor thetotal amount of radioactivity to which they are exposed. Eachmonth the badges are collected and scrutinized. Thus, em-ployees whose work places them in danger of excessive radio-active exposure are equipped with an “early warning system.”

Electrons

The positive charges in the nucleus of an atom are counter-balanced by negatively charged electrons orbiting at vary-ing distances around the nucleus. Thus, atoms with thesame number of protons and electrons are electrically neu-tral, having no net charge.

Electrons are maintained in their orbits by their attrac-tion to the positively charged nucleus. Sometimes otherforces overcome this attraction and an atom loses one ormore electrons. In other cases, atoms may gain additionalelectrons. Atoms in which the number of electrons doesnot equal the number of protons are known as ions, andthey carry a net electrical charge. An atom that has moreprotons than electrons has a net positive charge and iscalled a cation. For example, an atom of sodium (Na) thathas lost one electron becomes a sodium ion (Na+), with acharge of +1. An atom that has fewer protons than elec-trons carries a net negative charge and is called an anion. Achlorine atom (Cl) that has gained one electron becomes achloride ion (Cl–), with a charge of –1.

An atom consists of a nucleus of protons and neutronssurrounded by a cloud of electrons. The number of itselectrons largely determines the chemical properties ofan atom. Atoms that have the same number of protonsbut different numbers of neutrons are called isotopes.Isotopes of an atom differ in atomic mass but havesimilar chemical properties.

Chapter 2 The Nature of Molecules 21

Carbon-126 Protons6 Neutrons6 Electrons



FIGURE 2.3The three most abundantisotopes of carbon. Isotopesof a particular atom havedifferent numbers ofneutrons.

Electrons Determine the ChemicalBehavior of AtomsThe key to the chemical behavior of an atom lies in the ar-rangement of its electrons in their orbits. It is convenient tovisualize individual electrons as following discrete circularorbits around a central nucleus, as in the Bohr model of theatom. However, such a simple picture is not realistic. It isnot possible to precisely locate the position of any individualelectron precisely at any given time. In fact, a particularelectron can be anywhere at a given instant, from close tothe nucleus to infinitely far away from it.

However, a particular electron is more likely to be locat-ed in some positions than in others. The area around a nu-cleus where an electron is most likely to be found is calledthe orbital of that electron (figure 2.4). Some electron or-bitals near the nucleus are spherical (s orbitals), while oth-ers are dumbbell-shaped (p orbitals). Still other orbitals,more distant from the nucleus, may have different shapes.Regardless of its shape, no orbital may contain more thantwo electrons.

Almost all of the volume of an atom is empty space, be-cause the electrons are quite far from the nucleus relativeto its size. If the nucleus of an atom were the size of an ap-ple, the orbit of the nearest electron would be more than1600 meters away. Consequently, the nuclei of two atomsnever come close enough in nature to interact with eachother. It is for this reason that an atom’s electrons, not itsprotons or neutrons, determine its chemical behavior. This

also explains why the isotopes of an element, all of whichhave the same arrangement of electrons, behave the sameway chemically.

Energy within the Atom

All atoms possess energy, defined as the ability to do work.Because electrons are attracted to the positively chargednucleus, it takes work to keep them in orbit, just as it takeswork to hold a grapefruit in your hand against the pull ofgravity. The grapefruit is said to possess potential energy,the ability to do work, because of its position; if you wereto release it, the grapefruit would fall and its energy wouldbe reduced. Conversely, if you were to move the grapefruitto the top of a building, you would increase its potentialenergy. Similarly, electrons have potential energy of posi-tion. To oppose the attraction of the nucleus and move theelectron to a more distant orbital requires an input of en-ergy and results in an electron with greater potential ener-gy. This is how chlorophyll captures energy from lightduring photosynthesis (chapter 10)—the light excites elec-trons in the chlorophyll. Moving an electron closer to thenucleus has the opposite effect: energy is released, usuallyas heat, and the electron ends up with less potential energy(figure 2.5).

A given atom can possess only certain discrete amountsof energy. Like the potential energy of a grapefruit on a stepof a staircase, the potential energy contributed by the posi-tion of an electron in an atom can have only certain values.


1s Orbital

x x

y

z

Orbital for energy level K:one spherical orbital (1s)

2s Orbital

2p Orbitals

Composite ofall p orbitals

Orbitals for energy level L:one spherical orbital (2s) andthree dumbbell-shaped orbitals (2p)

z

y

FIGURE 2.4Electron orbitals. The lowest energy level or electron shell, which is nearest the nucleus, is level K. It is occupied by a single s orbital,referred to as 1s. The next highest energy level, L, is occupied by four orbitals: one s orbital (referred to as the 2s orbital) and three porbitals (each referred to as a 2p orbital). The four L-level orbitals compactly fill the space around the nucleus, like two pyramids set base-to-base.

Every atom exhibits a ladder of potential energy values,rather than a continuous spectrum of possibilities, a discreteset of orbits at particular distances from the nucleus.

During some chemical reactions, electrons are trans-ferred from one atom to another. In such reactions, the lossof an electron is called oxidation, and the gain of an elec-tron is called reduction (figure 2.6). It is important to real-ize that when an electron is transferred in this way, it keepsits energy of position. In organisms, chemical energy isstored in high-energy electrons that are transferred fromone atom to another in reactions involving oxidation andreduction.

Because the amount of energy an electron possesses isrelated to its distance from the nucleus, electrons that arethe same distance from the nucleus have the same energy,even if they occupy different orbitals. Such electrons aresaid to occupy the same energy level. In a schematic dia-gram of an atom (figure 2.7), the nucleus is represented as asmall circle and the electron energy levels are drawn as con-centric rings, with the energy level increasing with distancefrom the nucleus. Be careful not to confuse energy levels,which are drawn as rings to indicate an electron’s energy,with orbitals, which have a variety of three-dimensionalshapes and indicate an electron’s most likely location.

Electrons orbit a nucleus in paths called orbitals. Noorbital can contain more than two electrons, but manyorbitals may be the same distance from the nucleus and,thus, contain electrons of the same energy.


Energy released

Energylevel

3

Energylevel

2

Energylevel

1

–

M L K

Energylevel

1

Ener

gy a

bsor

bed

Energylevel

2

Energylevel

3

+

–

+ +++

+ +MLK

FIGURE 2.5Atomic energy levels. When an electronabsorbs energy, it moves to higher energylevels farther from the nucleus. When anelectron releases energy, it falls to lowerenergy levels closer to the nucleus.

FIGURE 2.6Oxidation and reduction. Oxidation is the loss of an electron;reduction is the gain of an electron.

Oxidation Reduction

�

–

�

Helium Nitrogen

7�

7n2�

2n KK L

Nucleus

L

M

N

K

Energy level

FIGURE 2.7Electron energy levels for helium and nitrogen. Gold ballsrepresent the electrons. Each concentric circle represents adifferent distance from the nucleus and, thus, a different electronenergy level.


1

H

1

H

3

Li4

Be

19

K

12

Mg

93

Np94

Pu95

Am96

Cm97

Bk98

Cf99

Es100

Fm101

Md102

No103

Lr

37

Rb38

Sr39

Y42

Mo45

Rh46

Pd47

Ag48

Cd49

In50

Sn51

Sb52

Te53

I54

Xe

21

Sc40

Zr

22

Ti23

V24

Cr25

Mn27

Co28

Ni29

Cu30

Zn36

Kr

5

B6

C6 C

8

O 2

He

55

Cs56

Ba72

Hf73

Ta74

W75

Re76

Os77

Ir78

Pt79

Au80

Hg81

Tl82

Pb83

Bi84

Po85

At86

Rn87

Fr88

Ra

57

La89

Ac104 105 106 107 108 109

58

Ce59

Pr60

Nd61

Pm62

Sm63

Eu64

Gd65

Tb66

Dy67

Ho68

Er69

Tm70

Yb71

Lu

90

Th91

Pa92

U

(Lanthanide series)

(Actinide series)

11

Na20

Ca

41

Nb43

Tc44

Ru

26

Fe

13

Al

31

Ga32

Ge

14

Si

7

N15

P

33

As

16

S

35

Br34

Se

9

F

18

Ar

10

Ne17

Cl

110

FIGURE 2.8Periodic table of the elements. In this representation, the frequency of elements that occur in the earth’s crust is indicated by the heightof the block. Elements found in significant amounts in living organisms are shaded in blue.

Kinds of AtomsThere are 92 naturally occurring elements, each with a dif-ferent number of protons and a different arrangement ofelectrons. When the nineteenth-century Russian chemistDmitri Mendeleev arranged the known elements in a tableaccording to their atomic mass (figure 2.8), he discoveredone of the great generalizations in all of science. Mendeleevfound that the elements in the table exhibited a pattern ofchemical properties that repeated itself in groups of eight el-ements. This periodically repeating pattern lent the table itsname: the periodic table of elements.

The Periodic Table

The eight-element periodicity that Mendeleev found isbased on the interactions of the electrons in the outer en-ergy levels of the different elements. These electrons arecalled valence electrons and their interactions are thebasis for the differing chemical properties of the elements.For most of the atoms important to life, an outer energy

level can contain no more than eight electrons; the chemi-cal behavior of an element reflects how many of the eightpositions are filled. Elements possessing all eight elec-trons in their outer energy level (two for helium) areinert, or nonreactive; they include helium (He), neon(Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon(Rn). In sharp contrast, elements with seven electrons (onefewer than the maximum number of eight) in their outerenergy level, such as fluorine (F), chlorine (Cl), andbromine (Br), are highly reactive. They tend to gain theextra electron needed to fill the energy level. Elementswith only one electron in their outer energy level, such aslithium (Li), sodium (Na), and potassium (K), are alsovery reactive; they tend to lose the single electron in theirouter level.

Mendeleev’s periodic table thus leads to a useful generali-zation, the octet rule (Latin octo, “eight”) or rule of eight:atoms tend to establish completely full outer energy levels.Most chemical behavior can be predicted quite accuratelyfrom this simple rule, combined with the tendency of at-oms to balance positive and negative charges.


Distribution of the Elements

Of the 92 naturally occurring elements on earth, only 11 arefound in organisms in more than trace amounts (0.01% orhigher). These 11 elements have atomic numbers less than21 and, thus, have low atomic masses. Table 2.1 lists thelevels of various elements in the human body; their levels inother organisms are similar. Inspection of this table suggeststhat the distribution of elements in living systems is by nomeans accidental. The most common elements inside or-ganisms are not the elements that are most abundant in the

earth’s crust. For example, silicon, aluminum, and iron con-stitute 39.2% of the earth’s crust, but they exist in traceamounts in the human body. On the other hand, carbon at-oms make up 18.5% of the human body but only 0.03% ofthe earth’s crust.

Ninety-two elements occur naturally on earth; onlyeleven of them are found in significant amounts in livingorganisms. Four of them—oxygen, hydrogen, carbon,nitrogen—constitute 96.3% of the weight of your body.


Table 2.1 The Most Common Elements on Earth and Their Distribution in the Human Body

Approximate Percent of Percent of Earth’s Crust Human Body

Element Symbol Atomic Number by Weight by Weight Importance or Function

Oxygen

SiliconAluminumIron

Calcium

Sodium

Potassium

Magnesium

Hydrogen

ManganeseFluorinePhosphorus

CarbonSulfurChlorineVanadiumChromiumCopperNitrogen

BoronCobaltZincSeleniumMolybdenumTinIodine

O

SiAlFe

Ca

Na

K

Mg

H

MnFP

CSClVCrCuN

BCoZnSeMoSnI

8

141326

20

11

19

12

1

259

15

616172324297

5273034425053

46.6

27.76.55.0

3.6

2.8

2.6

2.1

0.14

0.10.070.07

0.030.030.010.010.010.01

Trace

TraceTraceTraceTraceTraceTraceTrace

65.0

TraceTraceTrace

1.5

0.2

0.4

0.1

9.5

TraceTrace1.0

18.50.30.2

TraceTraceTrace3.3

TraceTraceTraceTraceTraceTraceTrace

Required for cellular respiration;component of water

Critical component of hemoglobin inthe bloodComponent of bones and teeth; trig-gers muscle contractionPrincipal positive ion outside cells;important in nerve functionPrincipal positive ion inside cells; im-portant in nerve functionCritical component of many energy-transferring enzymesElectron carrier; component of waterand most organic molecules

Backbone of nucleic acids; importantin energy transferBackbone of organic moleculesComponent of most proteinsPrincipal negative ion outside cells

Key component of many enzymesComponent of all proteins and nucleicacids

Key component of some enzymes

Key component of many enzymes

Component of thyroid hormone


Na

Sodium atom

Sodium ion

Chlorine atom

Chloride ion+ –

Na+

Cl

Cl�

(a)

FIGURE 2.9The formation of ionic bonds by sodium chloride. (a) When a sodium atom donates an electron to a chlorine atom, the sodium atombecomes a positively charged sodium ion, and the chlorine atom becomes a negatively charged chloride ion. (b) Sodium chloride forms ahighly regular lattice of alternating sodium ions and chloride ions.

NaCl crystal

Cl�

Cl�

Cl�Cl�

Cl�

Na�Na�

Na�

Na�

(b)

Ionic Bonds Form CrystalsA group of atoms held together by energy in a stable associ-ation is called a molecule. When a molecule contains atomsof more than one element, it is called a compound. Theatoms in a molecule are joined by chemical bonds; thesebonds can result when atoms with opposite charges attract(ionic bonds), when two atoms share one or more pairs ofelectrons (covalent bonds), or when atoms interact in otherways. We will start by examining ionic bonds, which formwhen atoms with opposite electrical charges (ions) attract.

A Closer Look at Table Salt

Common table salt, sodium chloride (NaCl), is a lattice ofions in which the atoms are held together by ionic bonds(figure 2.9). Sodium has 11 electrons: 2 in the inner energylevel, 8 in the next level, and 1 in the outer (valence) level.The valence electron is unpaired (free) and has a strong ten-dency to join with another electron. A stable configurationcan be achieved if the valence electron is lost to anotheratom that also has an unpaired electron. The loss of thiselectron results in the formation of a positively chargedsodium ion, Na+.

The chlorine atom has 17 electrons: 2 in the inner energylevel, 8 in the next level, and 7 in the outer level. Hence, oneof the orbitals in the outer energy level has an unpairedelectron. The addition of another electron to the outer levelfills that level and causes a negatively charged chloride ion,Cl–, to form.

When placed together, metallic sodium and gaseouschlorine react swiftly and explosively, as the sodium atomsdonate electrons to chlorine, forming Na+ and Cl– ions. Be-cause opposite charges attract, the Na+ and Cl– remain asso-ciated in an ionic compound, NaCl, which is electricallyneutral. However, the electrical attractive force holdingNaCl together is not directed specifically between particularNa+ and Cl– ions, and no discrete sodium chloride mole-cules form. Instead, the force exists between any one ion andall neighboring ions of the opposite charge, and the ions ag-gregate in a crystal matrix with a precise geometry. Such ag-gregations are what we know as salt crystals. If a salt such asNaCl is placed in water, the electrical attraction of the watermolecules, for reasons we will point out later in this chapter,disrupts the forces holding the ions in their crystal matrix,causing the salt to dissolve into a roughly equal mixture offree Na+ and Cl– ions.

An ionic bond is an attraction between ions of oppositecharge in an ionic compound. Such bonds are notformed between particular ions in the compound;rather, they exist between an ion and all of theoppositely charged ions in its immediate vicinity.


Covalent Bonds BuildStable MoleculesCovalent bonds form when two atomsshare one or more pairs of valenceelectrons. Consider hydrogen (H) as anexample. Each hydrogen atom has anunpaired electron and an unfilled outerenergy level; for these reasons the hy-drogen atom is unstable. When twohydrogen atoms are close to eachother, however, each atom’s electroncan orbit both nuclei. In effect, the nu-clei are able to share their electrons.The result is a diatomic (two-atom)molecule of hydrogen gas (figure 2.10).

The molecule formed by the two hy-drogen atoms is stable for three reasons:

1. It has no net charge. The di-atomic molecule formed as a resultof this sharing of electrons is notcharged, because it still containstwo protons and two electrons.

2. The octet rule is satisfied.Each of the two hydrogen atomscan be considered to have two or-biting electrons in its outer energylevel. This satisfies the octet rule,because each shared electron orbitsboth nuclei and is included in theouter energy level of both atoms.

3. It has no free electrons. Thebonds between the two atomsalso pair the two free electrons.

Unlike ionic bonds, covalent bondsare formed between two specific atoms, giving rise to true,discrete molecules. While ionic bonds can form regular crys-tals, the more specific associations made possible by covalentbonds allow the formation of complex molecular structures.

Covalent Bonds Can Be Very Strong

The strength of a covalent bond depends on the number ofshared electrons. Thus double bonds, which satisfy the oc-tet rule by allowing two atoms to share two pairs of elec-trons, are stronger than single bonds, in which only oneelectron pair is shared. This means more chemical energy isrequired to break a double bond than a single bond. Thestrongest covalent bonds are triple bonds, such as thosethat link the two nitrogen atoms of nitrogen gas molecules.Covalent bonds are represented in chemical formulations aslines connecting atomic symbols, where each line betweentwo bonded atoms represents the sharing of one pair ofelectrons. The structural formulas of hydrogen gas andoxygen gas are H—H and O�O, respectively, while theirmolecular formulas are H2 and O2.

Molecules with Several CovalentBonds

Molecules often consist of more thantwo atoms. One reason that larger mole-cules may be formed is that a given atomis able to share electrons with more thanone other atom. An atom that requirestwo, three, or four additional electronsto fill its outer energy level completelymay acquire them by sharing its elec-trons with two or more other atoms.

For example, the carbon atom (C)contains six electrons, four of which arein its outer energy level. To satisfy theoctet rule, a carbon atom must gain ac-cess to four additional electrons; that is,it must form four covalent bonds. Be-cause four covalent bonds may form inmany ways, carbon atoms are found inmany different kinds of molecules.

Chemical Reactions

The formation and breaking of chemi-cal bonds, the essence of chemistry, iscalled a chemical reaction. All chemi-cal reactions involve the shifting of at-oms from one molecule or ionic com-pound to another, without any changein the number or identity of the atoms.For convenience, we refer to the origi-nal molecules before the reaction startsas reactants, and the molecules result-ing from the chemical reaction as prod-ucts. For example:

A — B + C — D → A — C + B + Dreactants products

The extent to which chemical reactions occur is influ-enced by several important factors:

1. Temperature. Heating up the reactants increasesthe rate of a reaction (as long as the temperature isn’tso high as to destroy the molecules).

2. Concentration of reactants and products. Reac-tions proceed more quickly when more reactants areavailable. An accumulation of products typicallyspeeds reactions in the reverse direction.

3. Catalysts. A catalyst is a substance that increases therate of a reaction. It doesn’t alter the reaction’s equi-librium between reactants and products, but it doesshorten the time needed to reach equilibrium, oftendramatically. In organisms, proteins called enzymescatalyze almost every chemical reaction.

A covalent bond is a stable chemical bond formed whentwo atoms share one or more pairs of electrons.


FIGURE 2.10Hydrogen gas. (a) Hydrogen gas is adiatomic molecule composed of twohydrogen atoms, each sharing its electronwith the other. (b) The flash of fire thatconsumed the Hindenburg occurred whenthe hydrogen gas that was used to inflate thedirigible combined explosively with oxygengas in the air to form water.

H2 (hydrogen gas)

Covalent bond

+ +

–

–

(a)

(b)

Chemistry of WaterOf all the molecules that are common on earth, only wa-ter exists as a liquid at the relatively low temperatures thatprevail on the earth’s surface, three-fourths of which iscovered by liquid water (figure 2.11). When life was origi-nating, water provided a medium in which other moleculescould move around and interact without being held inplace by strong covalent or ionic bonds. Life evolved as aresult of these interactions, and it is still inextricably tiedto water. Life began in water and evolved there for 3 bil-lion years before spreading to land. About two-thirds ofany organism’s body is composed of water, and no organ-ism can grow or reproduce in any but a water-rich envi-ronment. It is no accident that tropical rain forests arebursting with life, while dry deserts appear almost lifelessexcept when water becomes temporarily plentiful, such asafter a rainstorm.

The Atomic Structure of Water

Water has a simple atomic structure. It consists of an oxy-gen atom bound to two hydrogen atoms by two single cova-lent bonds (figure 2.12a). The resulting molecule is stable: itsatisfies the octet rule, has no unpaired electrons, and carriesno net electrical charge.

The single most outstanding chemical property of wa-ter is its ability to form weak chemical associations withonly 5 to 10% of the strength of covalent bonds. This

property, which derives directly from the structure of wa-ter, is responsible for much of the organization of livingchemistry.

The chemistry of life is water chemistry. The way inwhich life first evolved was determined in large part bythe chemical properties of the liquid water in whichthat evolution occurred.


FIGURE 2.11Water takes many forms. As a liquid, water fills our rivers and runs down over the land to the sea. (a) The iceberg on which the penguinsare holding their meeting was formed in Antarctica from a huge block of ice that broke away into the ocean water. (b) When water coolsbelow 0°C, it forms beautiful crystals, familiar to us as snow and ice. However, water is not always plentiful. (c) At Badwater, in DeathValley, California, there is no hint of water except for the broken patterns of dried mud.

(a) (b) (c)

FIGURE 2.12Water has a simple molecular structure. (a) Each molecule iscomposed of one oxygen atom and two hydrogen atoms. Theoxygen atom shares one electron with each hydrogen atom. (b)The greater electronegativity of the oxygen atom makes the watermolecule polar: water carries two partial negative charges (δ–) nearthe oxygen atom and two partial positive charges (δ+), one on eachhydrogen atom.

δ–

δ–

δ–

δ–

δ+

δ+

δ+

δ+

104.5°

OxygenHydrogen

Hydrogen

Bohr model Ball-and-stick model

H

H

8+8n

+

+O

(a) (b)


Water Atoms Act Like TinyMagnetsBoth the oxygen and the hydrogen atoms attractthe electrons they share in the covalent bonds ofa water molecule; this attraction is called elec-tronegativity. However, the oxygen atom ismore electronegative than the hydrogen atoms,so it attracts the electrons more strongly than dothe hydrogen atoms. As a result, the sharedelectrons in a water molecule are far morelikely to be found near the oxygen nucleus thannear the hydrogen nuclei. This stronger attrac-tion for electrons gives the oxygen atom twopartial negative charges (δ–), as though the elec-tron cloud were denser near the oxygen atomthan around the hydrogen atoms. Because thewater molecule as a whole is electrically neu-tral, each hydrogen atom carries a partial positive charge (δ+).The Greek letter delta (δ) signifies a partial charge, muchweaker than the full unit charge of an ion.

What would you expect the shape of a water molecule tobe? Each of water’s two covalent bonds has a partial chargeat each end, δ– at the oxygen end and δ+ at the hydrogen end.The most stable arrangement of these charges is a tetrahe-dron, in which the two negative and two positive charges areapproximately equidistant from one another (figure 2.12b).The oxygen atom lies at the center of the tetrahedron, thehydrogen atoms occupy two of the apexes, and the partialnegative charges occupy the other two apexes. This resultsin a bond angle of 104.5° between the two covalent oxygen-hydrogen bonds. (In a regular tetrahedron, the bond angleswould be 109.5°; in water, the partial negative charges occu-py more space than the hydrogen atoms, and, therefore, theycompress the oxygen-hydrogen bond angle slightly.)

The water molecule, thus, has distinct “ends,” each with apartial charge, like the two poles of a magnet. (These partial

charges are much less than the unit charges of ions, how-ever.) Molecules that exhibit charge separation are calledpolar molecules because of their magnet-like poles, andwater is one of the most polar molecules known. The polarityof water underlies its chemistry and the chemistry of life.

Polar molecules interact with one another, as the δ– ofone molecule is attracted to the δ+ of another. Because manyof these interactions involve hydrogen atoms, they arecalled hydrogen bonds (figure 2.13). Each hydrogen bondis individually very weak and transient, lasting on averageonly 100,000

1,000,000 second (10–11 sec). However, the cumula-

tive effects of large numbers of these bonds can be enor-mous. Water forms an abundance of hydrogen bonds,which are responsible for many of its important physicalproperties (table 2.2).

The water molecule is very polar, with ends that exhibitpartial positive and negative charges. Opposite chargesattract, forming weak linkages called hydrogen bonds.


Hydrogen atom

Hydrogen bond

An organic molecule

Oxygen atom

δ–

δ+

FIGURE 2.13Structure of a hydrogen bond.

Table 2.2 The Properties of Water

Property Explanation Example of Benefit to Life

Cohesion

High specific heat

High heat of vaporizationLower density of ice

High polarity

Hydrogen bonds hold water molecules together

Hydrogen bonds absorb heat when they break, and releaseheat when they form, minimizing temperature changesMany hydrogen bonds must be broken for water to evapo-rateWater molecules in an ice crystal are spaced relatively farapart because of hydrogen bonding

Polar water molecules are attracted to ions and polar com-pounds, making them soluble

Leaves pull water upward from theroots; seeds swell and germinateWater stabilizes the temperature oforganisms and the environmentEvaporation of water cools bodysurfacesBecause ice is less dense than water,lakes do not freeze solid

Many kinds of molecules can movefreely in cells, permitting a diversearray of chemical reactions

Water Clings to Polar MoleculesThe polarity of water causes it to be attracted to other polarmolecules. When the other molecules are also water, the at-traction is referred to as cohesion. When the other mole-cules are of a different substance, the attraction is called ad-hesion. It is because water is cohesive that it is a liquid, andnot a gas, at moderate temperatures.

The cohesion of liquid water is also responsible for itssurface tension. Small insects can walk on water (figure2.14) because at the air-water interface all of the hydrogenbonds in water face downward, causing the molecules of thewater surface to cling together. Water is adhesive to anysubstance with which it can form hydrogen bonds. That iswhy substances containing polar molecules get “wet” whenthey are immersed in water, while those that are composedof nonpolar molecules (such as oils) do not.

The attraction of water to substances like glass with sur-face electrical charges is responsible for capillary action: if aglass tube with a narrow diameter is lowered into a beakerof water, water will rise in the tube above the level of thewater in the beaker, because the adhesion of water to theglass surface, drawing it upward, is stronger than the forceof gravity, drawing it down. The narrower the tube, thegreater the electrostatic forces between the water and theglass, and the higher the water rises (figure 2.15).

Water Stores Heat

Water moderates temperature through two properties: itshigh specific heat and its high heat of vaporization. Thetemperature of any substance is a measure of how rapidlyits individual molecules are moving. Because of the manyhydrogen bonds that water molecules form with one anoth-er, a large input of thermal energy is required to breakthese bonds before the individual water molecules can be-gin moving about more freely and so have a higher temper-ature. Therefore, water is said to have a high specific heat,which is defined as the amount of heat that must be ab-sorbed or lost by 1 gram of a substance to change its tem-perature by 1 degree Celsius (°C). Specific heat measuresthe extent to which a substance resists changing its temper-ature when it absorbs or loses heat. Because polar substanc-es tend to form hydrogen bonds, and energy is needed tobreak these bonds, the more polar a substance is, the higheris its specific heat. The specific heat of water (1 calo-rie/gram/°C) is twice that of most carbon compounds andnine times that of iron. Only ammonia, which is morepolar than water and forms very strong hydrogen bonds,has a higher specific heat than water (1.23calories/gram/°C). Still, only 20% of the hydrogen bondsare broken as water heats from 0° to 100°C.

Because of its high specific heat, water heats up moreslowly than almost any other compound and holds its tem-perature longer when heat is no longer applied. This char-acteristic enables organisms, which have a high water con-

tent, to maintain a relatively constant internal temperature.The heat generated by the chemical reactions inside cellswould destroy the cells, if it were not for the high specificheat of the water within them.

A considerable amount of heat energy (586 calories) is re-quired to change 1 gram of liquid water into a gas. Hence,water also has a high heat of vaporization. Because thetransition of water from a liquid to a gas requires the inputof energy to break its many hydrogen bonds, the evapora-tion of water from a surface causes cooling of that surface.Many organisms dispose of excess body heat by evaporativecooling; for example, humans and many other vertebratessweat.

At low temperatures, water molecules are locked into acrystal-like lattice of hydrogen bonds, forming the solid wecall ice (figure 2.16). Interestingly, ice is less dense than liquidwater because the hydrogen bonds in ice space the watermolecules relatively far apart. This unusual feature enablesicebergs to float. Were it otherwise, ice would cover nearly allbodies of water, with only shallow surface melting annually.


FIGURE 2.14Cohesion. Some insects, such as this water strider, literally walkon water. In this photograph you can see the dimpling the insect’sfeet make on the water as its weight bears down on the surface.Because the surface tension of the water is greater than the forcethat one foot brings to bear, the strider glides atop the surface ofthe water rather than sinking.

FIGURE 2.15Capillary action. Capillary actioncauses the water within a narrow tubeto rise above the surrounding water;the adhesion of the water to the glasssurface, which draws water upward, isstronger than the force of gravity,which tends to draw it down. Thenarrower the tube, the greater thesurface area available for adhesion for agiven volume of water, and the higherthe water rises in the tube.

Water Is a Powerful Solvent

Water is an effective solvent because of its ability to formhydrogen bonds. Water molecules gather closely aroundany substance that bears an electrical charge, whether thatsubstance carries a full charge (ion) or a charge separation(polar molecule). For example, sucrose (table sugar) iscomposed of molecules that contain slightly polar hydroxyl(OH) groups. A sugar crystal dissolves rapidly in water be-cause water molecules can form hydrogen bonds with indi-vidual hydroxyl groups of the sucrose molecules. There-fore, sucrose is said to be soluble in water. Every time asucrose molecule dissociates or breaks away from the crys-tal, water molecules surround it in a cloud, forming a hy-dration shell and preventing it from associating with oth-er sucrose molecules. Hydration shells also form aroundions such as Na+ and Cl– (figure 2.17).

Water Organizes Nonpolar Molecules

Water molecules always tend to form the maximum possi-ble number of hydrogen bonds. When nonpolar moleculessuch as oils, which do not form hydrogen bonds, are placed

in water, the water molecules act to exclude them. Thenonpolar molecules are forced into association with one an-other, thus minimizing their disruption of the hydrogenbonding of water. In effect, they shrink from contact withwater and for this reason they are referred to as hydropho-bic (Greek hydros, “water” and phobos, “fearing”). In con-trast, polar molecules, which readily form hydrogen bondswith water, are said to be hydrophilic (“water-loving”).

The tendency of nonpolar molecules to aggregate in wa-ter is known as hydrophobic exclusion. By forcing the hy-drophobic portions of molecules together, water causesthese molecules to assume particular shapes. Different mo-lecular shapes have evolved by alteration of the locationand strength of nonpolar regions. As you will see, much ofthe evolution of life reflects changes in molecular shapethat can be induced in just this way.

Water molecules, which are very polar, cling to oneanother, so that it takes considerable energy to separatethem. Water also clings to other polar molecules,causing them to be soluble in water solution, but watertends to exclude nonpolar molecules.


Watermolecules

Stablehydrogen bondsUnstable hydrogen bonds

(a) Liquid water (b) Ice

FIGURE 2.16The role of hydrogen bonds in an ice crystal. (a) In liquidwater, hydrogen bonds are not stable and constantly break and re-form. (b) When water cools below 0°C, the hydrogen bonds aremore stable, and a regular crystalline structure forms in which thefour partial charges of one water molecule interact with theopposite charges of other water molecules. Because water forms acrystal latticework, ice is less dense than liquid water and floats. Ifit did not, inland bodies of water far from the earth’s equatormight never fully thaw.

Hydration shells

Watermolecules

Saltcrystal

Na�

Cl�

Cl�

Cl�

Cl�

Na�

Na�

Na�

FIGURE 2.17Why salt dissolves in water. When a crystal of table salt dissolvesin water, individual Na+ and Cl– ions break away from the saltlattice and become surrounded by water molecules. Watermolecules orient around Cl– ions so that their partial positive polesface toward the negative Cl– ion; water molecules surrounding Na+

ions orient in the opposite way, with their partial negative polesfacing the positive Na+ ion. Surrounded by hydration shells, Na+

and Cl– ions never reenter the salt lattice.

Water IonizesThe covalent bonds within a water molecule sometimesbreak spontaneously. In pure water at 25°C, only 1 out ofevery 550 million water molecules undergoes this process.When it happens, one of the protons (hydrogen atom nu-clei) dissociates from the molecule. Because the dissociatedproton lacks the negatively charged electron it was sharingin the covalent bond with oxygen, its own positive charge isno longer counterbalanced, and it becomes a positivelycharged ion, H+. The rest of the dissociated water molecule,which has retained the shared electron from the covalentbond, is negatively charged and forms a hydroxide ion(OH–). This process of spontaneous ion formation is calledionization:

H2O → OH– + H+

water hydroxide ion hydrogenion (proton)

At 25°C, a liter of water contains 10,0001

,000 (or 10–7) mole of H+ ions. (A mole is defined as the weight ingrams that corresponds to the summed atomic masses of allof the atoms in a molecule. In the case ofH+, the atomicmass is 1, and a mole of H+ ions would weigh 1 gram. Onemole of any substance always contains 6.02 × 1023 mole-cules of the substance.) Therefore, the molar concentra-tion of hydrogen ions (represented as [H+]) in pure water is10–7 mole/liter. Actually, the hydrogen ion usually associ-ates with another water molecule to form a hydronium(H3O+) ion.

pH

A more convenient way to express the hydrogen ion concen-tration of a solution is to use the pH scale (figure 2.18).This scale defines pH as the negative logarithm of the hy-drogen ion concentration in the solution:

pH = –log [H+]

Because the logarithm of the hydrogen ion concentration issimply the exponent of the molar concentration of H+, thepH equals the exponent times –1. Thus, pure water, with an[H+] of 10–7 mole/liter, has a pH of 7. Recall that for everyH+ ion formed when water dissociates, an OH– ion is alsoformed, meaning that the dissociation of water produces H+

and OH– in equal amounts. Therefore, a pH value of 7 indi-cates neutrality—a balance between H+ and OH–—on thepH scale.

Note that the pH scale is logarithmic, which means that adifference of 1 on the scale represents a tenfold change inhydrogen ion concentration. This means that a solutionwith a pH of 4 has 10 times the concentration of H+ than ispresent in one with a pH of 5.

Acids. Any substance that dissociates in water to increasethe concentration of H+ ions is called an acid. Acidic solu-

tions have pH values below 7. The stronger an acid is, themore H+ ions it produces and the lower its pH. For exam-ple, hydrochloric acid (HCl), which is abundant in yourstomach, ionizes completely in water. This means that 10–1

mole per liter of HCl will dissociate to form 10–1 mole perliter of H+ ions, giving the solution a pH of 1. The pH ofchampagne, which bubbles because of the carbonic aciddissolved in it, is about 2.

Bases. A substance that combines with H+ ions when dis-solved in water is called a base. By combining with H+ ions,a base lowers the H+ ion concentration in the solution.Basic (or alkaline) solutions, therefore, have pH valuesabove 7. Very strong bases, such as sodium hydroxide(NaOH), have pH values of 12 or more.


10 �1

H+ IonConcentration

Examples ofSolutions

Hydrochloric acidStomach acid

1

pH Value

10 �2 Lemon juice2

10 �3 Vinegar, cola, beer3

10 �4 Tomatoes4

10 �5 Black coffeeNormal rainwater

5

10 �6 UrineSaliva

6

10 �7 Pure waterBlood

7

10 �8 Seawater8

10 �9 Baking soda9

10 �10 Great Salt Lake10

10 �11 Household ammonia11

10 �12

Household bleach12

10 �13

Oven cleaner13

10 �14 Sodium hydroxide14

FIGURE 2.18The pH scale. The pH value of a solution indicates itsconcentration of hydrogen ions. Solutions with a pH less than 7are acidic, while those with a pH greater than 7 are basic. Thescale is logarithmic, so that a pH change of 1 means a tenfoldchange in the concentration of hydrogen ions. Thus, lemon juiceis 100 times more acidic than tomato juice, and seawater is 10times more basic than pure water, which has a pH of 7.

Buffers

The pH inside almost all living cells, and in the fluid sur-rounding cells in multicellular organisms, is fairly close to 7.Most of the biological catalysts (enzymes) in living systemsare extremely sensitive to pH; often even a small change inpH will alter their shape, thereby disrupting their activitiesand rendering them useless. For this reason it is importantthat a cell maintain a constant pH level.

Yet the chemical reactions of life constantly produce acidsand bases within cells. Furthermore, many animals eat sub-stances that are acidic or basic; cola, for example, is a strong(although dilute) acidic solution. Despite such variations inthe concentrations of H+ and OH–, the pH of an organism iskept at a relatively constant level by buffers (figure 2.19).

A buffer is a substance that acts as a reservoir for hy-drogen ions, donating them to the solution when their con-centration falls and taking them from the solution whentheir concentration rises. What sort of substance will act inthis way? Within organisms, most buffers consist of pairs ofsubstances, one an acid and the other a base. The key bufferin human blood is an acid-base pair consisting of carbonicacid (acid) and bicarbonate (base). These two substances in-teract in a pair of reversible reactions. First, carbon dioxide(CO2) and H2O join to form carbonic acid (H2CO3), whichin a second reaction dissociates to yield bicarbonate ion(HCO3

–) and H+ (figure 2.20). If some acid or other sub-stance adds H+ ions to the blood, the HCO3

– ions act as abase and remove the excess H+ ions by forming H2CO3.Similarly, if a basic substance removes H+ ions from theblood, H2CO3 dissociates, releasing more H+ ions into theblood. The forward and reverse reactions that interconvertH2CO3 and HCO3

– thus stabilize the blood’s pH.The reaction of carbon dioxide and water to form car-

bonic acid is important because it permits carbon, essentialto life, to enter water from the air. As we will discuss inchapter 4, biologists believe that life first evolved in theearly oceans. These oceans were rich in carbon because ofthe reaction of carbon dioxide with water.

In a condition called blood acidosis, human blood, whichnormally has a pH of about 7.4, drops 0.2 to 0.4 points on thepH scale. This condition is fatal if not treated immediately.The reverse condition, blood alkalosis, involves an increase inblood pH of a similar magnitude and is just as serious.

The pH of a solution is the negative logarithm of the H+

ion concentration in the solution. Thus, low pH valuesindicate high H+ concentrations (acidic solutions), andhigh pH values indicate low H+ concentrations (basicsolutions). Even small changes in pH can be harmful tolife.


100

1

2

3

4

5

6

7

8

9

3Amount of base added

Bufferingrange

pH

4 52

FIGURE 2.19Buffers minimize changes in pH. Adding a base to a solutionneutralizes some of the acid present, and so raises the pH. Thus,as the curve moves to the right, reflecting more and more base, italso rises to higher pH values. What a buffer does is to make thecurve rise or fall very slowly over a portion of the pH scale, calledthe “buffering range” of that buffer.

FIGURE 2.20Buffer formation. Carbon dioxide and water combine chemically to form carbonic acid (H2CO3). The acid then dissociates in water,freeing H+ ions. This reaction makes carbonated beverages acidic, and produced the carbon-rich early oceans that cradled life.

H2O

Water

CO2

Carbon dioxide

H2CO3

Carbonic acid

HCO �3Bicarbonate

ion

H�

Hydrogenion

–

++

+

• The smallest stable particles of matter are protons, neutrons,and electrons, which associate to form atoms.

• The core, or nucleus, of an atom consists of protons andneutrons; the electrons orbit around the nucleus in a cloud.The farther an electron is from the nucleus, the faster itmoves and the more energy it possesses.

• The chemical behavior of an atom is largely determined bythe distribution of its electrons and in particular by the num-ber of electrons in its outermost (highest) energy level.There is a strong tendency for atoms to have a completelyfilled outer level; electrons are lost, gained, or shared untilthis condition is reached.



Chapter 2 Summary Questions Media Resources


1. An atom of nitrogen has 7protons and 7 neutrons. What isits atomic number? What is itsatomic mass? How many elec-trons does it have?2. How do the isotopes of a sin-gle element differ from eachother?3. The half-life of radium-226 is1620 years. If a sample of mate-rial contains 16 milligrams of ra-dium-226, how much will it con-tain in 1620 years? How muchwill it contain in 3240 years?How long will it take for thesample to contain 1 milligram ofradium-226?

• More than 95% of the weight of an organism consistsof oxygen, hydrogen, carbon, and nitrogen, all ofwhich form strong covalent bonds with one another.

• Ionic bonds form when electrons transfer from oneatom to another, and the resulting oppositely chargedions attract one another.

• Covalent bonds form when two atoms share elec-trons. They are responsible for the formation of mostbiologically important molecules.

4. What is the octet rule, andhow does it affect the chemicalbehavior of atoms?

5. What is the difference be-tween an ionic bond and a cova-lent bond? Give an example ofeach.



• The chemistry of life is the chemistry of water (H2O).The central oxygen atom in water attracts the elec-trons it shares with the two hydrogen atoms. Thischarge separation makes water a polar molecule.

• A hydrogen bond is formed between the partial posi-tive charge of a hydrogen atom in one molecule andthe partial negative charge of another atom, either inanother molecule or in a different portion of the samemolecule.

• Water is cohesive and adhesive, has a great capacityfor storing heat, is a good solvent for other polarmolecules, and tends to exclude nonpolar molecules.

• The H+ concentration in a solution is expressed bythe pH scale, in which pH equals the negative loga-rithm of the H+ concentration.

6. What types of atoms partici-pate in the formation of hydro-gen bonds? How do hydrogenbonds contribute to water’s highspecific heat?7. What types of molecules arehydrophobic? What types arehydrophilic? Why do these twotypes of molecules behave differ-ently in water?8. What is the pH of a solutionthat has a hydrogen ion concen-tration of 10–3 mole/liter?Would such a solution be acidicor basic?

• Atomic Structure

• Basic Chemistry• Atoms

• Bonds• Ionic Bonds

• Bonds

• Water• ph Scale

http://www.mhhe.com/raven6e http://www.biocourse.com

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