The macromolecules of most interest to molecular biologists—nucleic acids and proteins—are formed by the covalent joining of their constituent atoms. Because covalent bonds are relatively strong, stable, and not subject to spontaneous breakage under physiological conditions, they were once thought to be solely responsible for holding together the atoms in molecules. In contrast, weak chemical interactions, sometimes called weak chemical bonds, involve greater distances between atoms, are easily broken, and, individually, are transient. These properties, however, can be useful in biological systems, where transient chemical interactions are an essential part of cellular functions. For example, weak bonds mediate the interactions of proteins with small molecules, DNA, and/or other proteins. Such weak bonds enable countless critical processes in every cell, such as reversible binding of the hormone insulin to its receptor protein to regulate blood glucose levels or the interactions between adrenaline and receptor proteins that trigger the fight or-flight response.

When arranged in ordered groups, weak bonds can persist for a long time and can thus play central roles in the formation and stability of the active three-dimensional shapes of macromolecules. DNA is a case in point: although DNA is a linear polymer of covalently linked nucleotides, its shape and its ability to encode genetic information are determined by the stable yet dynamic double-helical structure that it adopts. This structure is defined by a large number of individually weak contacts between nucleotides that are not covalently bonded. Likewise, protein structures are largely determined by weak interactions between amino acid residues that are not necessarily adjacent in the polypeptide sequence. Therefore, although they are not strong enough individually to effectively bind two atoms together, weak chemical interactions play central roles in the structure and behavior of biological macromolecules.

Three kinds of weak chemical interactions are important in biological systems: van der Waals forces, hydrophobic interactions, and hydrogen bonds. Most macromolecules also use weak ionic interactions, along with these weak chemical interactions, to bind other molecules and to form three-dimensional structures.

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The Dutch chemist Johannes van der Waals was the first to document the intermolecular forces that result from the polarization of atoms. When two atoms approach each other, as they get closer, induced fluctuating charges between them cause a weak, nonspecific attractive interaction. This van der Waals interaction depends heavily on the distance between the interacting atoms. As the distance decreases below a certain point, a more powerful van der Waals repulsive force is caused by overlap of the atoms’ outer electron shells. The van der Waals radius of an atom is defined as the distance at which these attractive and repulsive forces are balanced and is characteristic for each atom (Figure 3-16).

The van der Waals bonding energy between two atoms that are separated by the sum of their van der Waals radii increases with the size of the atoms, but on average is only about 1 kcal/mol—just slightly above the average thermal energy of molecules at room temperature. This means that van der Waals forces are an effective binding force under physiological conditions only when they involve several atoms in each of the two interacting molecules. For several atoms to interact effectively in this way, the intermolecular fit must be exact, because the distance between any two interacting atoms must not be much different from the sum of their van der Waals radii.

The strongest kind of van der Waals contact arises when a macromolecule contains a surface that precisely fits the shape of the molecule that it binds. This is the case for antibodies, proteins that recognize antigens—specific molecules of viruses, bacteria, or other foreign particles that enter the body. Antibodies contain clefts with the same shape as the antigen they bind, enabling van der Waals contacts along the length of the bound antigen (Figure 3-17a). The additive effects of many van der Waals forces can be exceptionally strong; for instance, they are responsible for a gecko’s ability to climb vertically on a glass surface and hang there by a single toe (Figure 3-17b).

The hydrophobic effect arises from the strong tendency of water to exclude nonpolar groups, forcing these groups to aggregate in contact with one another. The word hydrophobic means “water-fearing,” which describes the apparent behavior of nonpolar molecules in water. For example, when drops of oil are added to water, they combine to form a larger drop. This happens because water molecules are attracted to one another, due to their polarity, whereas the nonpolar oil molecules have no charged regions to repel or attract other molecules. The attractive forces between water molecules result in an unfavorable organization of water molecules in the vicinity of hydrophobic molecules, decreasing the disorder, or entropy (see Section 3.6), of the nearby aqueous environment. As the oil drops aggregate, minimizing the surface area in contact with water, the net disorder of the surrounding water increases. Nonpolar molecules are sometimes said to undergo “hydrophobic interactions,” but their proximity is largely enforced by the effects of entropy in the surrounding aqueous medium. Such hydrophobic effects are common in biological molecules. For example, they are generally the dominant factor in stabilizing protein structures, and the energy required to unfold proteins goes mainly toward disrupting the stabilizing hydrophobic effect. The hydrophobic effect is critical to many other cellular functions as well, including insertion of protein molecules into membranes and secretion of hormones and other signaling molecules.

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The helical structures in DNA and RNA are stabilized by electronic interactions that arise from the stacked arrangement of base pairs. In accordance with the molecular orbital model introduced earlier, the electrons of the atoms in a base’s aromatic ring(s) are found in a decentralized cloud above and below the ring(s). Pi stacking is defined as the attractive, noncovalent interactions resulting from overlap between electrons of neighboring aromatic rings. This overlap of pi electrons of adjacent stacked nucleotides contributes to the stability of the nucleic acid’s helical structure (Figure 3-18).

A hydrogen bond is an attractive intermolecular force between two partial electrical charges of opposite polarity. As the name implies, one partner in the bond is a hydrogen, which must be covalently bonded to a strongly electronegative atom such as oxygen, nitrogen, or fluorine; this hydrogen is the hydrogen-bond donor. The electronegative atom attracts the electron cloud from around the hydrogen nucleus and, by decentralizing the cloud, leaves the hydrogen atom with a partial positive charge. This partial charge represents a large charge density that can attract the lone pair of electrons on another, nonhydrogen atom, which becomes the hydrogen-bond acceptor (Figure 3-19). Although other types of atoms can similarly acquire a partial positive charge when bonded to an electronegative element, only hydrogen is small enough to approach another atom or molecule close enough to undergo an energetically significant interaction.

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The hydrogen bond is not like a simple attraction between positive and negative charges at two points in space, but instead has directional preference and some characteristics of a covalent bond. This covalent character is more pronounced when acceptors hydrogen-bond with donors on more-electronegative atoms. For this reason, hydrogen bonds can vary in strength from very weak (a bonding energy of 2 kcal/mol) to fairly strong (7 kcal/mol).

The chemical properties of water are mostly due to hydrogen bonds that cause water molecules to have strong attractions for one another, but hydrogen bonds can and do form among many other kinds of molecules as well. In proteins, hydrogen bonds occur between the backbone carbonyl and imino groups to stabilize α helices and β sheets, the basic motifs of protein structure that we discuss in Chapter 4. In nucleic acids, hydrogen bonding between complementary pairs of nucleotide bases links one DNA strand to its partner, forming the double helix (see Figure 6-11).

The noncovalent interactions we have discussed (van der Waals forces, hydrophobic effects, pi bonds/stacking, and hydrogen bonds) are substantially weaker than covalent bonds. Just 1 kcal is needed to disrupt a mole of typical van der Waals interactions, but nearly 100 times more energy is required to break an equivalent number of covalent C–C or C–H bonds. The hydrophobic effect, too, is much more easily disrupted than covalent bonds, although it can be significantly strengthened in the presence of polar solvents such as concentrated salt solutions. Hydrogen bonds vary in strength depending on the polarity of the solvent and the alignment of the hydrogen-bonded atoms, but again, they are always weaker than covalent bonds. In aqueous solvent at 25°C, the available thermal energy is typically of the same order of magnitude as the strength of these weak interactions. Furthermore, the interaction of solute and solvent (water) molecules is nearly as favorable as any solute-solute interactions. Consequently, van der Waals forces, hydrophobic effects, and hydrogen bonds continually break and reform under physiological conditions.

As mentioned earlier, although these types of interactions are individually weak relative to covalent bonds, the cumulative effect of many such interactions can be significant. Macromolecules, including DNA, RNA, and proteins, contain so many sites of possible van der Waals or hydrophobic contacts and hydrogen bonding that the combined effect of these small binding forces is largely responsible for their molecular structure. For macromolecules, the most stable structure is usually that in which weak interactions are maximized. This principle determines the folding of a single polypeptide or polynucleotide chain into its three-dimensional shape. Complete unfolding of the structure requires the removal of all these interactions at the same time. And because these contacts are breaking and re-forming rapidly and randomly, such synchronized disruptions are very unlikely. The molecular stability conferred by many weak interactions is therefore much greater than one might expect intuitively, based on a simple summation of many small binding energies.

A special class of weak interactions in macromolecules involves the water molecules that are invariably bound to surface and interior sites by hydrogen bonds. Sometimes these water molecules are so well positioned that they behave as though they are an integral part of the macromolecule. In many cases, bound water molecules are essential to macromolecular function. Certain DNA-binding proteins, for example, use water molecules that are integral to their structure to help recognize specific DNA sequences. In RNA structures, water molecules bridge nucleotide bases that are involved in nontraditional base pairing and in interactions involving three nucleotides (base triples), thereby enabling unique three-dimensional structures to form (Figure 3-20).

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Interactions among DNA, RNA, and protein molecules, and interactions between these macromolecules and small organic molecules, are also mediated by weak, noncovalent chemical interactions. Such contacts involving information-carrying macromolecules govern how and when cells replicate, repair and recombine their DNA, synthesize RNA and proteins, detect and respond to chemical signals, and conduct all the other activities essential for life.

For example, the noncovalent binding of a protein to another protein, to a small molecule (such as a hormone), or to another macromolecule (such as a nucleic acid) may involve several hydrogen bonds and one or more ionic, hydrophobic, and/or van der Waals interactions (Figure 3-21). These weak contacts contribute, overall, to an energetically favorable interaction. It is worth noting, however, that because each hydrogen bond between two groups in a protein or nucleic acid forms at the expense of two hydrogen bonds between the same groups and two water molecules, the net stabilization is not as great as it might seem. The large size of nucleic acids and proteins provides extensive molecular surfaces with many opportunities for weak interactions with other molecules. The energetic favorability of such interactions reflects the molecular surface complementarity and the resulting large numbers of weak interactions of polar, charged, and hydrophobic groups on the surfaces of these molecules.

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