Human Metabolism. Keith N. Frayn

of methane has the three-dimensional structure shown in Figure 1.3a. The outer electron ‘cloud’ has a very even distribution over the four hydrogen atoms, all of which have an equal tendency to pull electrons their way. The molecule has no distinct electrical poles – it is non-polar. Because of this very even distribution of electrons, molecules near each other have little tendency to interact. In contrast, in the water molecule (Figure 1.3b) the oxygen atom has a distinct tendency to pull electrons its way, shifting the distribution of the outer electron cloud so that it is more dense over the oxygen atom, and correspondingly less dense elsewhere. Therefore, the molecule has a rather negatively charged region around the central oxygen atom, and correspondingly positively charged regions around the hydrogen atoms. Thus, it has distinct electrical poles – it is a relatively polar molecule. It is easy to imagine that water molecules near to each other will interact. Like electrical charges repel each other, unlike charges attract. This gives water molecules a tendency to line up so that the positive regions of one attract the negative region of an adjacent molecule (Figure 1.3b). So, water molecules, unlike those of methane, tend to ‘stick together’: the energy needed to break them apart and form a gas is much greater than for methane, and hence water is a liquid while methane is a gas. The latent heat of evaporation of water is 2.5 kJ g⁻¹, whereas that of methane is 0.6 kJ g⁻¹. Note that the polarity of the water molecule is not as extreme as that of an ion – it is merely a rather uneven distribution of electrons, but enough to affect its properties considerably.

Figure 1.3 (a) Three-dimensional structure of the methane molecule and (b) the molecular structure of water. (a) The hydrogen atoms of methane (CH₄) are arranged symmetrically in space, at the corners of a tetrahedron. (b) The molecular structure of water. Top: view of the ‘electron cloud’ surrounding the molecule; bottom, interactions between water molecules. The molecule has a degree of polarity, and this leads to electrical interactions between neighbouring molecules by the formation of hydrogen bonds. These bonds are not strong compared with covalent bonds, and are constantly being formed and broken. Nevertheless, they provide sufficient attraction between the molecules to account for the fact that water is a liquid at room temperature whereas the non-polar methane is a gas.

      The contrast between water and methane may be extended to larger molecules. Organic compounds composed solely of carbon and hydrogen – for instance, the alkanes or ‘paraffins’ – all have the property of extreme non-polarity: the chemical (covalent) bond between carbon and hydrogen atoms leads to a very even distribution of electrons, and the molecules have little interaction with each other. A result is that polar molecules, such as those of water, and non-polar molecules, such as those of alkanes, do not mix well: the water molecules tend to bond to each other and to exclude the non-polar molecules, which can themselves pack together very closely because of the lack of interaction between them. In fact, there is an additional form of direct attraction between non-polar molecules, the van der Waals forces. Random fluctuations in the density of the electron cloud surrounding a molecule lead to minor, transient degrees of polarity; these induce an opposite change in a neighbouring molecule, with the result that there is a transient attraction between them. These are very weak attractions, however, and the effect of the exclusion by water is considerably stronger. The non-polar molecules are said to be hydrophobic (water fearing or water hating).

A strong contrast is provided by an inorganic ionic compound such as sodium chloride. The sodium and chlorine atoms in sodium chloride are completely ionised under almost all conditions. They pack very regularly in crystals in a cubic form. The strength of their attraction for each other means that considerable energy is needed to disrupt this regular packing – sodium chloride does not melt until heated above 800 °C. And yet it dissolves very readily in water – that is, the individual ions become separated from their close packing arrangement rather as they would on melting. Why? Because the water molecules, by virtue of their polarity, are able to come between the ions and reduce their attraction for each other. In fact, each of the charged sodium and chloride ions will become surrounded by a ‘shell’ of water molecules, shielding it from the attraction or repulsion of other ions. Sodium chloride is said to be hydrophilic – water loving. The terms polar and hydrophilic are for the most part interchangeable. Similarly, the terms non-polar and hydrophobic are virtually synonymous.

      Ionic compounds, the extreme examples of polarity, are not confined to inorganic chemistry. Organic molecules may include ionised groups. These may be almost entirely ionised under normal conditions – for instance, the esters of orthophosphoric acid (‘phosphate groups’), as in the compounds AMP, ADP, and ATP, in metabolites such as glucose 6-phosphate, and in phospholipids. Most of the organic acids involved in intermediary metabolism, such as lactic acid, pyruvic acid, and the long-chain carboxylic acids (fatty acids), are also largely ionised at physiological hydrogen ion concentrations (Box 1.1). Thus, generation of lactic acid during exercise raises the hydrogen ion concentration (the acidity) both within the cells where it is produced, and generally within the body, since it is released into the bloodstream.

Box 1.1 Ionisation state of some acids at normal hydrogen ion concentrations

      The normal pH in blood plasma is around 7.4. (It may be somewhat lower within cells, down to about 6.8.) This corresponds to a hydrogen ion concentration of 3.98 × 10⁻⁸ mol l⁻¹ (since – log₁₀ of 3.98 × 10⁻⁸ is 7.4).

      The equation for ionisation of an acid HA is:

      this equilibrium is described by the equation:

      where K_i is the dissociation or ionisation constant and is a measure of the strength of the acid: the higher the value of K_i the stronger (i.e. the more dissociated) the acid.

      K_i in the equation above relates the concentrations expressed in molar terms (e.g. mol/l). (Strictly, it is not the concentrations but the ‘effective ion concentrations’ or ion activities which are related; these are not quite the same as concentrations because of inter-ion attractions. In most biological systems, however, in which the concentrations are relatively low, it is a close approximation to use concentrations. If activities are used, then the symbol K_a is used for the dissociation constant of an acid.)

Some biological acids and their K_a values are listed in Table 1.1.1, together with a calculation of the proportion ionised at typical pH (7.4).

      The calculation is done as follows (using acetic acid as an example):