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Lodish H, Berk A, Zipursky SL, et al. Molecular Cell Biology. 4th edition. New York: W. H. Freeman; 2000.

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Covalent bonds, which hold the atoms within anindividual molecule together, are formed by the sharing of electrons in the outer atomicorbitals. The distribution of shared as well as unshared electrons in outer orbitals is a majordeterminant of the three-dimensional shape and chemical reactivity of molecules. For instance,as we learn in Chapter 3, the shape of proteins iscrucial to their function and their interactions with small molecules. In this section, wediscuss important properties of covalent bonds and describe the structure of carbohydrates toillustrate how the geometry of bonds determines the shape of small biological molecules.

Each Atom Can Make a Defined Number of Covalent Bonds

Electrons move around the nucleus of an atom in clouds called orbitals,which lie in a series of concentric shells, or energy levels; electrons inouter shells have more energy than those in inner shells. Each shell has a maximum number ofelectrons that it can hold. Electrons fill the innermost shells of an atom first; then theouter shells. The energy level of an atom is lowest when all of its orbitals are filled, and anatom’s reactivity depends on how many electrons it needs to complete its outermostorbital. In most cases, in order to fill the outermost orbital, the electrons within it formcovalent bonds with other atoms. A covalent bond thus holds two atoms close together becauseelectrons in their outermost orbitals are shared by both atoms.

Most of the molecules in living systems contain only six different atoms: hydrogen, carbon,nitrogen, phosphorus, oxygen, and sulfur. The outermost orbital of each atom has acharacteristic number of electrons:


These atoms readily form covalent bonds with other atoms and rarely exist as isolatedentities. As a rule, each type of atom forms a characteristic number of covalent bonds withother atoms.

For example, a hydrogen atom, with one electron in its outer shell, forms only one bond, suchthat its outermost orbital becomes filled with two electrons. A carbon atom has four electronsin its outermost orbitals; it usually forms four bonds, as in methane (CH4), inorder to fill its outermost orbital with eight electrons. The single bonds in methane thatconnect the carbon atom with each hydrogen atom contain two shared electrons, one donated fromthe C and the other from the H, and the outer (s) orbital of each H atom isfilled by the two shared electrons:


Nitrogen and phosphorus each have five electrons in their outer shells, which can hold up toeight electrons. Nitrogen atoms can form up to four covalent bonds. In ammonia(NH3), the nitrogen atom forms three covalent bonds; one pair of electrons aroundthe atom (the two dots on the right) are in an orbital not involved in a covalent bond:


In the ammonium ion (NH4+), the nitrogen atom forms fourcovalent bonds, again filling the outermost orbital with eight electrons:


Phosphorus can form up to five covalent bonds, as in phosphoric acid(H3PO4). The H3PO4 molecule is actually a“resonance hybrid,” a structure between the two forms shown below in whichnonbonding electrons are shown as pairs of dots:
In theresonance hybrid on the right, one of the electrons from the P=O double bond hasaccumulated around the O atom, giving it a net negative charge and leaving the P atom with anet positive charge. The resonance hybrid on the left, in which the P atom forms the maximumfive covalent bonds, has no charged atoms. Esters of phosphoric acid form the backbone ofnucleic acids, as discussed in Chapter 4;phosphates also play key roles in cellular energetics (Chapter 16) and in the regulation of cell function (Chapters 13 and 20).

The difference between the bonding patterns of nitrogen and phosphorus is primarily due tothe relative sizes of the two atoms: the smaller nitrogen atom has only enough space toaccommodate four bonding pairs of electrons around it without creating destructive repulsionsbetween them, whereas the larger sphere of the phosphorus atom allows more electron pairs to bearranged around it without the pairs being too close together.

Both oxygen and sulfur contain six electrons in their outermost orbitals. However, an atom ofoxygen usually forms only two covalent bonds, as in molecular oxygen, O2:

Primarily because its outermost orbital is larger than that of oxygen, sulfur can form as fewas two covalent bonds, as in hydrogen sulfide (H2S), or as many as six, as in sulfurtrioxide (SO3) or sulfuric acid (H2SO4):
Esters of sulfuric acid are important constituents of the proteoglycansthat compose part of the extracellular matrix surrounding most animal cells (Chapter 22).

The Making or Breaking of Covalent Bonds Involves Large Energy Changes

Covalent bonds tend to be very stable because the energies required to break or rearrangethem are much greater than the thermal energy available at room temperature (25 °C) orbody temperature (37 °C). For example, the thermal energy at 25 °C is less than1 kilocalorie per mole (kcal/mol), whereas the energy required to break a C—C bond inethane is about 83 kcal/mol:
whereΔH represents the difference in the total energy of all of thebonds (the enthalpy) in the reactants and in the products.*The positive value indicates that an input of energy is needed to cause the reaction, andthat the products contain more energy than the reactants. The high energy needed for breakageof the ethane bond means that at room temperature (25 °C) well under 1 in1012 ethane molecules exists as a pair of ·CH3 radicals. Thecovalent bonds in biological molecules have ΔH values similar tothat of the C—C bond in ethane (Table2-1).

Covalent Bonds Have Characteristic Geometries

When two or more atoms form covalent bonds with another central atom, these bonds areoriented at precise angles to one another. The angles are determined by the mutual repulsion ofthe outer electron orbitals of the central atom. These bond angles give each molecule itscharacteristic shape (Figure 2-2). In methane, forexample, the central carbon atom is bonded to four hydrogen atoms, whose positions define thefour points of a tetrahedron, so that the angle between any two bonds is 109.5°. Likemethane, the ammonium ion also has a tetrahedral shape. In these molecules, each bond is asingle bond, a single pair of electrons shared between two atoms. When twoatoms share two pairs of electrons — for example, when a carbonatom is linked to only three other atoms — the bond is adouble bond:
In thiscase, the carbon atom and all three atoms linked to it lie in the same plane (Figure 2-3). Atoms connected by a double bond cannot rotatefreely about the bond axis, while those in a single bond generally can. The rigid planarityimposed by double bonds has enormous significance for the shape of large biological moleculessuch as proteins and nucleic acids. (In triple bonds, two atoms share sixelectrons. These are rare in biological molecules.)

Figure 2-2

Bond angles give these water and methane molecules their distinctive shapes. Each molecule is represented in three ways. The atoms in the ball-and-stick models aresmaller than they actually are in relation to bond length, to show the bond angles clearly.The (more...)

Figure 2-3

In an ethylene molecule, the carbon atoms are connected by a double bond, causing allthe atoms to lie in the same plane. Unlike atoms connected by a single bond, which usually can rotate freely about the bondaxis, those connected by a double bond cannot. (more...)

All outer electron orbitals, whether or not they are involved in covalent bond formation,contribute to the properties of a molecule, in particular to its shape. For example, the outershell of the oxygen atom in a water molecule has two pairs of nonbonding electrons; the twopairs of electrons in the H—O bonds and the two pairs of nonbonding electrons form analmost perfect tetrahedron. However, the orbitals of the nonbonding electrons have a highelectron density and thus tend to repel each other, compressing the angle between the covalentH—O—H bonds to 104.5° rather than the 109.5° in atetrahedron (see Figure 2-2).

Electrons Are Shared Unequally in Polar Covalent Bonds

In a covalent bond, one or more pairs of electrons are shared between two atoms. In certaincases, the bonded atoms exert different attractions for the electrons of the bond, resulting inunequal sharing of the electrons. The power of an atom in a molecule to attract electrons toitself, called electronegativity, is measured on a scale from 4.0 (forfluorine, the most electronegative atom) to a hypothetical zero (Figure 2-4). Knowing the electronegativity of two atoms allows us to predictwhether a covalent bond can form between them; if the differences in electronegativity areconsiderable — as in sodium andchloride — an ionic bond, rather than a covalent bond, willform. This type of interaction is discussed in a later section.

Figure 2-4

Electronegativity values of main-group elements in the periodic table. Atoms located to the upper right tend to have high electronegativity, fluorine being themost electronegative. Elements with low electronegativity values, such as the metalslithium, (more...)

In a covalent bond in which the atoms either are identical or have the sameelectronegativity, the bonding electrons are shared equally. Such a bond is said to be nonpolar. This is the case for C—C andC—H bonds. However, if two atoms differ in electronegativity, the bond is said to bepolar. One end of a polar bond has a partialnegative charge (δ−), and the other end has a partial positivecharge (δ+). In an O—H bond, for example, the oxygenatom, with an electronegativity of 3.4, attracts the bonded electrons more than does thehydrogen atom, which has an electronegativity of 2.2. As a result, the bonding electrons spendmore time around the oxygen atom than around the hydrogen. Thus the O—H bondpossesses an electric dipole, a positive charge separated from an equal butopposite negative charge. We can think of the oxygen atom of the O—H bond as having,on average, a charge of 25 percent of an electron, with the H atom having an equivalentpositive charge. The dipole moment of the O—H bond is a function ofthe size of the positive or negative charge and the distance separating the charges.

In a water molecule both hydrogen atoms are on the same side of the oxygen atom. As a result,the side of the molecule with the two H atoms has a slight net positive charge, whereas theother side has a slight net negative charge. Because of this separation of positive andnegative charges, the entire molecule has a net dipole moment (Figure 2-5). Some molecules, such as the linear molecule CO2, have twopolar bonds:
Because the dipole moments of the two C=Obonds point in opposite directions, they cancel each other out, resulting in a molecule withouta net dipole moment.

Figure 2-5

The water molecule has two polar O—H bonds and a net dipole moment. The symbol δ represents a partial charge (a weaker charge than the one on anelectron or a proton), and each of the polar H—O bonds has a dipole moment. Thenet (more...)

Asymmetric Carbon Atoms Are Present in Most Biological Molecules

A carbon (or any other) atom bonded to four dissimilar atoms or groups is said to beasymmetric. The bonds formed by an asymmetric carbonatom can be arranged in threedimensional space in two different ways, producingmolecules that are mirror images of each other. Such molecules are called opticalisomers, or stereoisomers. One isomer issaid to be right-handed and the other left-handed, a property calledchirality. Most molecules in cells contain at least one asymmetric carbon atom, often called a chiral carbon atom. The different stereoisomers of amolecule usually have completely different biological activities.

Amino Acids

Except for glycine, all amino acids, the building blocks of the proteins, have one chiralcarbon atom, called the α carbon, orCα, which is bonded to four different atoms or groupsof atoms. In the amino acid alanine, for instance, this carbon atom is bonded to—NH2, —COOH, —H, and —CH3(Figure 2-6). By convention, the two mirror-imagestructures are called the D (dextro) and the L (levo)isomers of the amino acid. The two isomers cannot be interconverted without breaking achemical bond. With rare exceptions, only the L forms of amino acids are found in proteins. Wediscuss the properties of amino acids and the covalent peptide bond that links them into longchains in Chapter 3.

Figure 2-6

Stereoisomers of the amino acid alanine. The asymmetric α carbon is black. Although the chemical properties of suchoptical isomers are identical, their biological activities are distinct.


The three-dimensional structures of carbohydrates provide another excellent example of thestructural and biological importance of chiral carbon atoms, even in simple molecules. Acarbohydrate is constructed of carbon (carbo-) plus hydrogen and oxygen(-hydrate, or water). The formula for the simplestcarbohydrates — the monosaccharides, or simple sugars — is(CH2O)n, where n equals 3, 4, 5, 6, or 7. All monosaccharides contain hydroxyl(—OH) groups and either an aldehyde or a keto group:

In the linear form of D-glucose (C6H12O6),the principal source of energy for most cells in higher organisms, carbon atoms 2, 3, 4, and 5are asymmetric (Figure 2-7, top). Ifthe hydrogen atom and the hydroxyl group attached to carbon atom 2 (C2) wereinterchanged, the resulting molecule would be a different sugar, D-mannose, and could not beconverted to glucose without breaking and making covalent bonds. Enzymes can distinguishbetween this single point of difference.

Figure 2-7

Three alternative configurations of D-glucose. The ring forms, shown as Haworth projections, are generated from the linear molecule byreaction of the aldehyde at carbon 1 with the hydroxyl on carbon 5 or carbon 4.

D-Glucose can exist in three different forms: a linear structure and two differenthemiacetal ring structures (see Figure 2-7). If thealdehyde group on carbon 1 reacts with the hydroxyl group on carbon 5, the resultinghemiacetal, D-glucopyranose, contains a six-member ring. Similarly, condensation of thehydroxyl group on carbon 4 with the aldehyde group results in the formation ofD-glucofuranose, a hemiacetal containing a five-member ring. Although all three forms ofD-glucose exist in biological systems, the pyranose form is by far the most abundant.

The planar depiction of the pyranose ring shown in Figure2-7 is called a Haworth projection. When a linear molecule ofD-glucose forms a pyranose ring, carbon 1 becomes asymmetric, so two stereoisomers (calledanomers) of D-glucopyranose are possible. The hydroxyl group attached tocarbon 1 “points” down (below the plane of projection) inα-D-glucopyranose, as shown in Figure 2-7,and points up (above the plane of projection) in the β anomer. In aqueous solutionthe α and β anomers readily interconvert spontaneously; at equilibriumthere is about one-third α anomer and two-thirds β, with very little of theopen-chain form. Because enzymes can distinguish between the α and βanomers of D-glucose, these forms have specific biological roles.

Most biologically important sugars are six-carbon sugars, or hexoses, that are structurally related to D-glucose. Mannose, as noted, isidentical with glucose except for the orientation of the substituents on carbon 2. In Haworthprojections of the pyranose forms of glucose and mannose, the hydroxyl group on carbon 2 ofglucose points downward, whereas that on mannose points upward (Figure 2-8). Similarly, galactose, another hexose, differs from glucoseonly in the orientation of the hydroxyl group on carbon 4.

Figure 2-8

Haworth projections of the structures of glucose, mannose, and galactose in theirpyranose forms. The hydroxyl groups with different orientations from those of glucose arehighlighted.

The Haworth projection is an oversimplification be-cause the actual pyranose ring is notplanar. Rather, sugar molecules adopt a conformation in which each of the ring carbons is atthe center of a tetrahedron, just like the carbon in methane (see Figure 2-2). The preferred conformation of pyranose structures is the chair(Figure 2-9). In this conformation, the bonds goingfrom a ring carbon to nonring atoms may take two directions: axial (perpendicular to the ring)and equatorial (in the plane of the ring).

Figure 2-9

Chair conformations of glucose, mannose, and galactose in their pyranoseforms. The chair is the most stable conformation of a six-membered ring. (In an alternativeform, called the boat, both carbon 1 and carbon 4 lie above the plane ofthe ring.) The (more...)

The L isomers of sugars are virtually unknown in biological systems except for L-fucose. Oneof the unsolved mysteries of molecular evolution is why only D isomers of sugars and L isomersof amino acids were utilized, and not the chemically equivalent L sugars and D aminoacids.

α and β Glycosidic Bonds Link Monosaccharides

In addition to the monosaccharides discussed above, two common disaccharides, lactose and sucrose, occur naturally (Figure 2-10). A disaccharide consists of two monosaccharides linked togetherby a C—O—C bridge called a glycosidicbond. The disaccharide lactose is the major sugar in milk; sucrose is a principalproduct of plant photosynthesis and is refined into common table sugar.

Figure 2-10

The formation of glycosidic linkages generate the disaccharides lactose andsucrose. The lactose linkage is β(1 → 4); the sucroselinkage is α(1 → 2). In any glycosidic linkage,carbon 1 (more...)

In the formation of any glycosidic bond, the carbon 1 atom of one sugar molecule reacts witha hydroxyl group of another. As in the formation of most biopolymers, the linkage isaccompanied by the loss of water. In principle, a large number of different glycosidic bondscan be formed between two sugar residues. Glucose could be bonded to fructose, for example, byany of the following linkages: α(1 → 1),α(1 → 2),α(1 → 3),α(1 → 4),α(1 → 6),β(1 → 1),β(1 → 2),β(1 → 3),β(1 → 4), orβ(1 → 6), where α or β specifiesthe conformation at carbon 1 in glucose and the number following the arrow indicates thefructose carbon to which the glucose is bound. Only theα(1 → 2) linkage occurs in sucrose because of thespecificity of the enzyme (the biological catalyst) for the linking reaction.

Glycosidic linkages also join chains of monosaccharides into longer polymers, called polysaccharides, some of which function as reservoirsfor glucose. The most common storage carbohydrate in animal cells is glycogen, a very long, highly branched polymer of glucose units linkedtogether mainly by α(1 → 4) glycosidic bonds. Asmuch as 10 percent by weight of the liver can be glycogen. The primary storage carbohydrate inplant cells, starch, also is a glucose polymerwith α(1 → 4) linkages. It occurs in two forms,amylose, which is unbranched, and amylopectin, which has some branches. In contrast to glycogenand starch, some polysaccharides, such as cellulose, have structural and other nonstorage functions. An unbranched polymer ofglucose linked together by β(1 → 4) glycosidicbonds, cellulose is the major constituent of plant cell walls and is the most abundant organicchemical on earth. Because of the different linkages between the glucose units, cellulose formslong rods, whereas glycogen and starch form coiled helices. Human digestive enzymes canhydrolyze α(1 → 4) glycosidic bonds, but notβ(1 → 4) bonds, between glucose units; for thisreason humans can digest starch but not cellulose. The synthesis and utilization of thesepolysaccharides are described in later chapters.

 Covalent bonds, which bind the atoms composing a molecule ina fixed orientation, consist of pairs of electrons shared by two atoms. Relatively highenergies are required to break them (50 – 200 kcal/mol).
 Most molecules in cells contain at least one chiral(asymmetric) carbon atom, which is bonded to four dissimilar atoms. Such molecules can existas optical isomers, designated D and L, which have identical chemical properties butcompletely different biological activities. In biological systems, nearly all amino acids areL isomers and nearly all sugars are D isomers.

A calorie is defined as the amount of thermal energy required to heat 1 cm3 ofwater by 1 °C from 14 °C to 15 °C. Many biochemistry textbooks usethe joule (J), but the two units can be interconverted quite readily (1cal = 4.184 J). The energy changes in chemical reactions,such as the making or breaking of chemical bonds, are measured in kilocalories per mole inthis book (1 kcal = 1000 cal). One mole of any substance isthe amount that contains 6.02 × 1023 items of thatsubstance, which is known as Avogadro’s number. Thus, one canspeak of a mole of photons, or 6.02 × 1023photons. The weight of a mole of a substance in grams (g) is the same as its molecularweight. For example, the molecular weight of water is 18, so the weight of 1 mole of water,or 6.02 × 1023 water molecules, is 18 g.

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