1.2.1 The biologically significant molecular structure of water
In our biological view of thermodynamics, we defined a set of standard biological conditions that are different from those usually encountered in classical discussions on chemistry or physics. One of the more prominently different set of conditions we noted were those applied to water, which is the universal biological solvent. Thus, it seems appropriate to turn our attention to water and consider the basis for its special treatment in biological chemistry.
First, water is the dominant molecular species in most animals, plants, and cells comprising about 65% of the mass of the human body. Water is the main biological solvent, with most of the biochemical reactions in microbes, plants, and animals taking place in the aqueous compartments that make up cells and complex organisms. The special properties of water that make it the universal biological solvent are related to the spatial distribution of electrons in the two covalent bonds that exist between the oxygen atoms and the hydrogen atoms that comprise molecular water (Fig. 1.2). The nucleus of the oxygen atom with its eight protons is strongly electrophylic compared to the hydrogen s-
Figure 1.2 Dipolar water molecule. The figure shows the two covalent H-O bonds and the partial charges on the hydrogen and oxygen atoms which are the result of bonding electrons being displaced toward the oxygen nucleus with its relatively high positive charge.
atom with its single positive nuclear proton. As a consequence of these unequal nuclear charges, the electron pairs that constitute the covalent bonds are, on average, markedly displaced toward the oxygen atom and generate a partial negative charge (S—) on the oxygen as illustrated in Fig. 1. 2. The displacement of the bonding electrons away from the hydrogen atoms asymmetrically unshields the protons that comprise the hydrogen nuclei. It results in the appearance of a partial positive charge (S+) on each of the hydrogen atoms, and that partial charge is directed away from the oxygen atom. Because of these partial charge separations, water molecules are often described as molecular dipoles having partial charges on opposite sides of the molecule.
In bulk water, the charged regions of individual water molecules interact electrostatically with oppositely charged regions of neighboring molecules resulting in a highly structured matrix of molecules where partial positive charges on one molecule interact with partial negative charges on neighboring molecule, as shown in Fig. 1.3. The resulting dipole-dipole interaction is unusually strong having some of the qualities of both electrostatic and covalent bonds, and thus has acquired a special name, the hydrogen bond (H bond). The enthalpic contribution to the energy of a system comprised of hydrogen bonds is in the range of 1 to 10 kcal/mol of H bond. The energy of H bonds in water is generally being taken to be about 5 kcal/mol. The latter value, although about 10 times lower than the covalent O-H bonds in water molecules, imparts considerable stability to aqueous biological systems since the room temperature (25°C) energy available to thermally destabilize bonds is about 10 times lower than the water-water H bond energy.
In the crystalline form of water that we deal with in everyday life, it is generally concluded that each water molecule is hydrogen bonded to four other water molecules resulting in formation of the familiar material known as ice. As the temperature of water ice is raised above the
Figure 1.3 Tetrahedrally bonded crystalline ice. The central water molecule is shaded. The broken lines represent hydrogen bonds.
freezing point, hydrogen bonds break resulting in the formation of evanescent clusters of hydrogen-bonded molecules that wink in and out of existence as H bonds form and reform with a half-life of about 10 picoseconds (ps). The loss of the stable crystalline structure results in liquid water having a higher density than ice-state water and provides the basis for the fact that ice, with its lower density, always floats in water. The continuing presence of considerable intermolecular H bond structure in room temperature water and its practical significance is easily evidenced by the inexpert high divers who learn very rapidly not to have large areas of their body, like their belly, flop into the water, fast and all at the same time.
In biological systems, where water is the solvent and solutes are present in relatively low concentration, the water is generally considered to be present at a concentration that is not significantly different than that of pure water (55.5 M). This high concentration of water is very effective in dissolving and solvating ions such as sodium (Na+) and chloride (Cl-). In these organized complexes, water dipoles form highly oriented spherical shells around each individual ion with the countercharge of the water dipoles oriented toward the charge of the solvated ion. The spatial extent and stability of these oriented hydration spheres are directly proportional to the surface charge density of the ion. Likewise, polar organic molecules such as alcohols and sugars, as well as ionized organic molecules such as acids and bases, dissolve in water as a consequence of similar hydration effects. As we will see later in our examination of biological membranes (Sec. 1.6.2), the interactions of water dipoles is also important in helping stabilize the organization of nonpolar molecules like fats and lipids into organized lipid structures like biological membranes.
While pure water and water in biological systems are largely in the form of rapidly shifting molecular clusters, some water dissociates producing equivalent amounts of hydrogen (H ) and hydroxyl (OH ) ions. The extent of this dissociation is small but significant, resulting in equal H+ and OH- concentrations of 10-7. The special importance of hydrogen ions in biology and chemistry has led to the development of the pH scale of H+ measurement to facilitate the discussion and description of H concentrations. The pH of a hydrogen ion containing solution is simply the negative log of the hydrogen ion concentration as shown in Eq. (1.13).
Thus, the pH of a 10 7 M solution of H+ is 7.0. Although there are striking individual exceptions, it is generally the case that the intracellular and extracellular compartments associated with living cells and organisms are tightly regulated in the range between pH 5.0 and pH 7.5, and hydrogen ion concentrations outside this range are unfavorable for the vast majority of life forms. One biologically important exception to this generalization is the pH of the stomach, in which the gastric juice has a pH of about 0.7 to 3.8 (see Clinical Box 1.1).
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