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Chapter: Biochemistry: Water: The Solvent for Biochemical Reactions

Hydrogen Bonds

Hydrogen Bonds
In addition to the interactions discussed, another important type of noncovalent interaction exists: hydrogen bonding.

Hydrogen Bonds

In addition to the interactions discussed, another important type of noncovalent interaction exists: hydrogen bonding. Hydrogen bonding is of electrostatic origin and can be considered a special case of dipole–dipole interaction. When hydrogen is covalently bonded to a very electronegative atom such as oxygen or nitrogen, it has a partial positive charge due to the polar bond, a situation that does not occur when hydrogen is covalently bonded to carbon. This partial positive charge on hydrogen can interact with an unshared (nonbonding) pair of electrons (a source of negative charge) on another electronegative atom. All three atoms lie in a straight line, forming a hydrogen bond. This arrangement allows for the greatest possible partial positive charge on the hydrogen and, consequently, for the strongest possible interaction with the unshared pair of electrons on the second electronegative atom (Figure 2.6). The group comprising the electronegative atom that is covalently bonded to hydrogen is called the hydrogen-bond donor, and the electronegative atom that contributes the unshared pair of electrons to the interaction is the hydrogen-bondacceptor. The hydrogen is not covalently bonded to the acceptor in the usualdescription of hydrogen bonding.

Recent research has cast some doubt on this view, with experimental evi-dence to indicate some covalent character in the hydrogen bond.

Why does water have such interesting and unique properties?

A consideration of the hydrogen-bonding sites in HF, H2O, and NH3 can yield some useful insights. Figure 2.7 shows that water constitutes an optimum situ-ation in terms of the number of hydrogen bonds that each molecule can form. Water has two hydrogens to enter into hydrogen bonds and two unshared pairs of electrons on the oxygen to which other water molecules can be hydrogen-bonded. 

Each water molecule is involved in four hydrogen bonds-as a donor in two and as an acceptor in two. Hydrogen fluoride has only one hydrogen to enter into a hydrogen bond as a donor, but it has three unshared pairs of electrons on the fluorine that could bond to other hydrogens. Ammonia has three hydrogens to donate to a hydrogen bond but only one unshared pair of electrons, on the nitrogen.

The geometric arrangement of hydrogen-bonded water molecules has important implications for the properties of water as a solvent. The bond angle in water is 104.3°, as was shown in Figure 2.1, and the angle between the unshared pairs of electrons is similar. The result is a tetrahedral arrange-ment of water molecules. Liquid water consists of hydrogen-bonded arrays that resemble ice crystals; each of these arrays can contain up to 100 water molecules. The hydrogen bonding between water molecules can be seen more clearly in the regular lattice structure of the ice crystal (Figure 2.8). There are several differences, however, between hydrogen-bonded arrays of this type in liquid water and the structure of ice crystals. In liquid water, hydrogen bonds are constantly breaking and new ones are constantly forming, with some mol-ecules breaking off and others joining the cluster. A cluster can break up and re-form in 10210 to 10211 seconds in water at 25°C. An ice crystal, in contrast, has a more-or-less-stable arrangement of hydrogen bonds, and of course its number of molecules is many orders of magnitude greater than 100.

Hydrogen bonds are much weaker than normal covalent bonds. Whereas the energy required to break the O-H covalent bond is 460 kJ mol21 (110 kcal mol21), the energy of hydrogen bonds in water is about 20 kJ mol21 (5 kcal mol21) (Table 2.3). Even this comparatively small amount of energy is enough to affect the properties of water drastically, especially its melting point, its boil-ing point, and its density relative to the density of ice. Both the melting point and the boiling point of water are significantly higher than would be predicted for a molecule of this size (Table 2.4). 

Other substances of about the same molecular weight, such as methane and ammonia, have much lower melting and boiling points. The forces of attraction between the molecules of these substances are weaker than the attraction between water molecules, because of the number and strength of their hydrogen bonds. The energy of this attraction must be overcome to melt ice or boil water.

Ice has a lower density than liquid water because the fully hydrogen bonded array in an ice crystal is less densely packed than that in liquid water. Liquid water is less extensively hydrogen-bonded and thus is denser than ice. Thus, ice cubes and icebergs float. Most substances contract when they freeze, but the opposite is true of water. 

In cold weather, the cooling systems of cars require antifreeze to prevent freezing and expansion of the water, which could crack the engine block. In laboratory procedures for cell fractionation, the same principle is used in a method of disrupting cells with several cycles of freezing and thawing. Finally, aquatic organisms can survive in cold climates because of the density difference between ice and liquid water; lakes and rivers freeze from top to bottom rather than vice versa.

Hydrogen bonding also plays a role in the behavior of water as a solvent. If a polar solute can serve as a donor or an acceptor of hydrogen bonds, not only can it form hydrogen bonds with water but it can also be involved in nonspe-cific dipole2dipole interactions. Figure 2.9 shows some examples. Alcohols, amines, carboxylic acids, and esters, as well as aldehydes and ketones, can all form hydrogen bonds with water, so they are soluble in water. It is difficult to overstate the importance of water to the existence of life on the Earth, and it is difficult to imagine life based on another solvent. The following Biochemical Connections box explores some of the implications of this statement.

Other Biologically Important Hydrogen Bonds

Hydrogen bonds have a vital involvement in stabilizing the three-dimensional structures of biologically important molecules, including DNA, RNA, and proteins. The hydrogen bonds between complementary bases are one of the most striking characteristics of the double-helical structure of DNA. Transfer RNA also has a complex three-dimensional structure characterized by hydrogen-bonded regions. Hydrogen bonding in proteins gives rise to two important structures, the α-helix and β-pleated sheet conformations. Both types of conformation are widely encountered in proteins. Table 2.5 summarizes some of the most important kinds of hydrogen bonds in biomolecules.


A hydrogen bond is a special example of a dipole–dipole bond.


Water molecules are extensively hydrogen bonded.


The ability to form strong hydrogen bonds is responsible for the many unique characteristics of water, such as its very high melting point and boiling point for a molecule of its size.


The three-dimensional structures of many important biomolecules, including proteins and nucleic acids, are stabilized by hydrogen bonds.


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