Molecular Cell Biology.

Section 2.1Covalent Bonds

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

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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 in outer shells have more energy than those in inner shells. Each shell has a maximum number of electrons that it can hold. Electrons fill the innermost shells of an atom first; then the outer shells. The energy level of an atom is lowest when all of its orbitals are filled, and an atom’s reactivity depends on how many electrons it needs to complete its outermost orbital. In most cases, in order to fill the outermost orbital, the electrons within it form covalent bonds with other atoms. A covalent bond thus holds two atoms close together because electrons 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 a characteristic number of electrons:

These atoms readily form covalent bonds with other atoms and rarely exist as isolated entities. As a rule, each type of atom forms a characteristic number of covalent bonds with other atoms.
For example, a hydrogen atom, with one electron in its outer shell, forms only one bond, such that its outermost orbital becomes filled with two electrons. A carbon atom has four electrons in its outermost orbitals; it usually forms four bonds, as in methane (CH₄), in order to fill its outermost orbital with eight electrons. The single bonds in methane that connect the carbon atom with each hydrogen atom contain two shared electrons, one donated from the C and the other from the H, and the outer (s) orbital of each H atom is filled by the two shared electrons:

Nitrogen and phosphorus each have five electrons in their outer shells, which can hold up to eight electrons. Nitrogen atoms can form up to four covalent bonds. In ammonia (NH₃), the nitrogen atom forms three covalent bonds; one pair of electrons around the 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 four covalent 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 which nonbonding electrons are shown as pairs of dots:

In the resonance hybrid on the right, one of the electrons from the P=O double bond has accumulated around the O atom, giving it a net negative charge and leaving the P atom with a net positive charge. The resonance hybrid on the left, in which the P atom forms the maximum five covalent bonds, has no charged atoms. Esters of phosphoric acid form the backbone of nucleic 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 to the relative sizes of the two atoms: the smaller nitrogen atom has only enough space to accommodate four bonding pairs of electrons around it without creating destructive repulsions between them, whereas the larger sphere of the phosphorus atom allows more electron pairs to be arranged around it without the pairs being too close together.
Both oxygen and sulfur contain six electrons in their outermost orbitals. However, an atom of oxygen 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 few as two covalent bonds, as in hydrogen sulfide (H₂S), or as many as six, as in sulfur trioxide (SO₃) or sulfuric acid (H₂SO₄):

Esters of sulfuric acid are important constituents of the proteoglycans that compose part of the extracellular matrix surrounding most animal cells (Chapter 22).

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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 rearrange them are much greater than the thermal energy available at room temperature (25 °C) or body temperature (37 °C). For example, the thermal energy at 25 °C is less than 1 kilocalorie per mole (kcal/mol), whereas the energy required to break a C—C bond in ethane is about 83 kcal/mol:

where ΔH represents the difference in the total energy of all of the bonds (the enthalpy) in the reactants and in the products.^*The positive value indicates that an input of energy is needed to cause the reaction, and that the products contain more energy than the reactants. The high energy needed for breakage of the ethane bond means that at room temperature (25 °C) well under 1 in 10¹² ethane molecules exists as a pair of ·CH₃ radicals. The covalent bonds in biological molecules have ΔH values similar to that of the C—C bond in ethane (Table 2-1).

Table 2-1

The Energy Required to Break Some Important Covalent Bonds Found in Biological Molecules*.

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Covalent Bonds Have Characteristic Geometries

When two or more atoms form covalent bonds with another central atom, these bonds are oriented at precise angles to one another. The angles are determined by the mutual repulsion of the outer electron orbitals of the central atom. These bond angles give each molecule its characteristic shape (Figure 2-2). In methane, for example, the central carbon atom is bonded to four hydrogen atoms, whose positions define the four points of a tetrahedron, so that the angle between any two bonds is 109.5°. Like methane, the ammonium ion also has a tetrahedral shape. In these molecules, each bond is a single bond, a single pair of electrons shared between two atoms. When two atoms share two pairs of electrons — for example, when a carbon atom is linked to only three other atoms — the bond is a double bond:

In this case, 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 rotate freely about the bond axis, while those in a single bond generally can. The rigid planarity imposed by double bonds has enormous significance for the shape of large biological molecules such as proteins and nucleic acids. (In triple bonds, two atoms share six electrons. 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 are smaller 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 all the atoms to lie in the same plane. Unlike atoms connected by a single bond, which usually can rotate freely about the bond axis, 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 outer shell of the oxygen atom in a water molecule has two pairs of nonbonding electrons; the two pairs of electrons in the H—O bonds and the two pairs of nonbonding electrons form an almost perfect tetrahedron. However, the orbitals of the nonbonding electrons have a high electron density and thus tend to repel each other, compressing the angle between the covalent H—O—H bonds to 104.5° rather than the 109.5° in a tetrahedron (see Figure 2-2).

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Electrons Are Shared Unequally in Polar Covalent Bonds

In a covalent bond, one or more pairs of electrons are shared between two atoms. In certain cases, the bonded atoms exert different attractions for the electrons of the bond, resulting in unequal sharing of the electrons. The power of an atom in a molecule to attract electrons to itself, called electronegativity, is measured on a scale from 4.0 (for fluorine, the most electronegative atom) to a hypothetical zero (Figure 2-4). Knowing the electronegativity of two atoms allows us to predict whether a covalent bond can form between them; if the differences in electronegativity are considerable — as in sodium and chloride — an ionic bond, rather than a covalent bond, will form. 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 the most electronegative. Elements with low electronegativity values, such as the metals lithium, (more...)

In a covalent bond in which the atoms either are identical or have the same electronegativity, the bonding electrons are shared equally. Such a bond is said to be nonpolar. This is the case for C—C and C—H bonds. However, if two atoms differ in electronegativity, the bond is said to be polar. One end of a polar bond has a partial negative charge (δ⁻), and the other end has a partial positive charge (δ⁺). In an O—H bond, for example, the oxygen atom, with an electronegativity of 3.4, attracts the bonded electrons more than does the hydrogen atom, which has an electronegativity of 2.2. As a result, the bonding electrons spend more time around the oxygen atom than around the hydrogen. Thus the O—H bond possesses an electric dipole, a positive charge separated from an equal but opposite 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 equivalent positive charge. The dipole moment of the O—H bond is a function of the 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 the other side has a slight net negative charge. Because of this separation of positive and negative charges, the entire molecule has a net dipole moment (Figure 2-5). Some molecules, such as the linear molecule CO₂, have two polar bonds:

Because the dipole moments of the two C=O bonds point in opposite directions, they cancel each other out, resulting in a molecule without a 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 an electron or a proton), and each of the polar H—O bonds has a dipole moment. The net (more...)

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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 be asymmetric. The bonds formed by an asymmetric carbon atom can be arranged in threedimensional space in two different ways, producing molecules that are mirror images of each other. Such molecules are called optical isomers, or stereoisomers. One isomer is said to be right-handed and the other left-handed, a property called chirality. Most molecules in cells contain at least one asymmetric carbon atom, often called a chiral carbon atom. The different stereoisomers of a molecule usually have completely different biological activities.

Amino Acids

Except for glycine, all amino acids, the building blocks of the proteins, have one chiral carbon atom, called the α carbon, or C_α, which is bonded to four different atoms or groups of 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-image structures are called the D (dextro) and the L (levo) isomers of the amino acid. The two isomers cannot be interconverted without breaking a chemical bond. With rare exceptions, only the L forms of amino acids are found in proteins. We discuss the properties of amino acids and the covalent peptide bond that links them into long chains in Chapter 3.

Figure 2-6

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

Carbohydrates

The three-dimensional structures of carbohydrates provide another excellent example of the structural and biological importance of chiral carbon atoms, even in simple molecules. A carbohydrate is constructed of carbon (carbo-) plus hydrogen and oxygen (-hydrate, or water). The formula for the simplest carbohydrates — 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 5 are asymmetric (Figure 2-7, top). If the hydrogen atom and the hydroxyl group attached to carbon atom 2 (C₂) were interchanged, the resulting molecule would be a different sugar, D-mannose, and could not be converted to glucose without breaking and making covalent bonds. Enzymes can distinguish between 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 by reaction 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 different hemiacetal ring structures (see Figure 2-7). If the aldehyde group on carbon 1 reacts with the hydroxyl group on carbon 5, the resulting hemiacetal, D-glucopyranose, contains a six-member ring. Similarly, condensation of the hydroxyl group on carbon 4 with the aldehyde group results in the formation of D-glucofuranose, a hemiacetal containing a five-member ring. Although all three forms of D-glucose exist in biological systems, the pyranose form is by far the most abundant.
The planar depiction of the pyranose ring shown in Figure 2-7 is called a Haworth projection. When a linear molecule of D-glucose forms a pyranose ring, carbon 1 becomes asymmetric, so two stereoisomers (called anomers) of D-glucopyranose are possible. The hydroxyl group attached to carbon 1 “points” down (below the plane of projection) in α-D-glucopyranose, as shown in Figure 2-7, and points up (above the plane of projection) in the β anomer. In aqueous solution the α and β anomers readily interconvert spontaneously; at equilibrium there is about one-third α anomer and two-thirds β, with very little of the open-chain form. Because enzymes can distinguish between the α and β anomers 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, is identical with glucose except for the orientation of the substituents on carbon 2. In Haworth projections of the pyranose forms of glucose and mannose, the hydroxyl group on carbon 2 of glucose points downward, whereas that on mannose points upward (Figure 2-8). Similarly, galactose, another hexose, differs from glucose only in the orientation of the hydroxyl group on carbon 4.

Figure 2-8

Haworth projections of the structures of glucose, mannose, and galactose in their pyranose forms. The hydroxyl groups with different orientations from those of glucose are highlighted.

The Haworth projection is an oversimplification be-cause the actual pyranose ring is not planar. Rather, sugar molecules adopt a conformation in which each of the ring carbons is at the 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 going from 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 pyranose forms. The chair is the most stable conformation of a six-membered ring. (In an alternative form, called the boat, both carbon 1 and carbon 4 lie above the plane of the ring.) The (more...)

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

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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 together by a C—O—C bridge called a glycosidic bond. The disaccharide lactose is the major sugar in milk; sucrose is a principal product of plant photosynthesis and is refined into common table sugar.

Figure 2-10

The formation of glycosidic linkages generate the disaccharides lactose and sucrose. The lactose linkage is β(1 → 4); the sucrose linkage 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 with a hydroxyl group of another. As in the formation of most biopolymers, the linkage is accompanied by the loss of water. In principle, a large number of different glycosidic bonds can be formed between two sugar residues. Glucose could be bonded to fructose, for example, by any 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), or β(1 → 6), where α or β specifies the conformation at carbon 1 in glucose and the number following the arrow indicates the fructose carbon to which the glucose is bound. Only the α(1 → 2) linkage occurs in sucrose because of the specificity 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 reservoirs for glucose. The most common storage carbohydrate in animal cells is glycogen, a very long, highly branched polymer of glucose units linked together mainly by α(1 → 4) glycosidic bonds. As much as 10 percent by weight of the liver can be glycogen. The primary storage carbohydrate in plant cells, starch, also is a glucose polymer with α(1 → 4) linkages. It occurs in two forms, amylose, which is unbranched, and amylopectin, which has some branches. In contrast to glycogen and starch, some polysaccharides, such as cellulose, have structural and other nonstorage functions. An unbranched polymer of glucose linked together by β(1 → 4) glycosidic bonds, cellulose is the major constituent of plant cell walls and is the most abundant organic chemical on earth. Because of the different linkages between the glucose units, cellulose forms long rods, whereas glycogen and starch form coiled helices. Human digestive enzymes can hydrolyze α(1 → 4) glycosidic bonds, but not β(1 → 4) bonds, between glucose units; for this reason humans can digest starch but not cellulose. The synthesis and utilization of these polysaccharides are described in later chapters.

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SUMMARY

•  Covalent bonds, which bind the atoms composing a molecule in a fixed orientation, consist of pairs of electrons shared by two atoms. Relatively high energies are required to break them (50 – 200 kcal/mol).
•  In covalent bonds between unlike atoms that differ in electronegativity, the bonding electrons are distributed unequally. In such polar bonds, one end has a partial positive charge and the other end has a partial negative charge (see Figure 2-5).
•  Most molecules in cells contain at least one chiral (asymmetric) carbon atom, which is bonded to four dissimilar atoms. Such molecules can exist as optical isomers, designated D and L, which have identical chemical properties but completely different biological activities. In biological systems, nearly all amino acids are L isomers and nearly all sugars are D isomers.
•  Glucose and other hexoses can exist in three forms: an open-chain linear structure, a six-member (pyranose) ring, and a five-member (furanose) ring (see Figure 2-7). In biological systems, the pyranose form of D- glucose predominates. The two possible stereoisomers of D-glucopyranose (the α and β anomers) differ in the orientation of the hydroxyl group attached to carbon 1.
•  Glycosidic bonds link carbon 1 of one monosaccharide to a hydroxyl group on another sugar, leading to formation of disaccharides and polysaccharides. Many different glycosidic bonds are theoretically possible between two sugar residues, but the enzymes that make and break these bonds are specific for the α or β anomer of one sugar and a particular hydroxyl group on the other.

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Footnotes

*

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