The reactions inside a cell occur in an aqueous environment. Life on Earth began in the ocean, and the conditions in that primeval environment put a permanent stamp on the chemistry of living things. Life therefore hinges on the chemical properties of water, which are reviewed in Panel 2–2, pp. 92–93.
In each water molecule (H2O) the two H atoms are linked to the O atom by covalent bonds. The two bonds are highly polar because the O is strongly attractive for electrons, whereas the H is only weakly attractive. Consequently, there is an unequal distribution of electrons in a water molecule, with a preponderance of positive charge on the two H atoms and of negative charge on the O. When a positively charged region of one water molecule (that is, one of its H atoms) approaches a negatively charged region (that is, the O) of a second water molecule, the electrical attraction between them can result in a hydrogen bond. These bonds are much weaker than covalent bonds and are easily broken by the random thermal motions that reflect the heat energy of the molecules. Thus, each bond lasts only a short time. But the combined effect of many weak bonds can be profound. For example, each water molecule can form hydrogen bonds through its two H atoms to two other water molecules, producing a network in which hydrogen bonds are being continually broken and formed. It is only because of the hydrogen bonds that link water molecules together that water is a liquid at room temperature—with a high boiling point and high surface tension—rather than a gas.
Molecules, such as alcohols, that contain polar bonds and that can form hydrogen bonds with water dissolve readily in water. Molecules carrying charges (ions) likewise interact favorably with water. Such molecules are termed hydrophilic, meaning that they are water-loving. Many of the molecules in the aqueous environment of a cell necessarily fall into this category, including sugars, DNA, RNA, and most proteins. Hydrophobic (water-hating) molecules, by contrast, are uncharged and form few or no hydrogen bonds, and so do not dissolve in water. Hydrocarbons are an important example. In these molecules all of the H atoms are covalently linked to C atoms by a largely nonpolar bond; thus they cannot form effective hydrogen bonds to other molecules (see Panel 2–1, p. 90). This makes the hydrocarbon as a whole hydrophobic—a property that is exploited in cells, whose membranes are constructed from molecules that have long hydrocarbon tails, as we see in Chapter 10.
four Types of noncovalent attractions help Bring molecules Together in Cells
Much of biology depends on the specific binding of different molecules caused by three types of noncovalent bonds: electrostatic attractions (ionic bonds), hydrogen bonds, and van der Waals attractions; and on a fourth factor that can push molecules together: the hydrophobic force. The properties of the four types of noncovalent attractions are presented in Panel 2–3 (pp. 94–95). Although each
Figure 2–2 Some energies important for cells. a crucial property of any bond— covalent or noncovalent—is its strength. Bond strength is measured by the amount of energy that must be supplied to break it, expressed in units of either kilojoules per mole (kJ/mole) or kilocalories per mole (kcal/mole). Thus if 100 kJ of energy must be supplied to break 6 × 1023 bonds of a specific type (that is, 1 mole of these bonds), then the strength of that bond is 100 kJ/mole. note that, in this diagram, energies are compared on a logarithmic scale. Typical strengths and lengths of the main classes of chemical bonds are given in Table 2–1.