Water is the most abundant substance in living systems, making up 70% or more of the weight of most organisms. Water pervades all portions of every cell and is the medium in which the transport of nutrients, the enzyme-catalyzed reactions of metabolism, and the transfer of chemical energy occur. The first living organisms probably arose in the primeval oceans; evolution was shaped by the properties of the medium in which it occurred. All aspects of cell structure and function are adapted to the physical and chemical properties of water. This chapter begins with descriptions of these physical and chemical properties. The strong attractive forces between water molecules result in water's solvent properties. The slight tendency of water to ionize is also of crucial importance to the structure and function of biomolecules, and we will review the topic of ionization in terms of equilibrium constants, pH, and titration curves. Finally, we will consider the way in which aqueous solutions of weak acids or bases and their salts act as buffers against pH changes in biological systems. The water molecule and its ionization products, H+ and OH -, profoundly influence the structure, self assembly, and properties of all cellular components, including enzymes and other proteins, nucleic acids, and lipids. The noncovalent interactions responsible for the specificity of "recognition" among biomolecules are decisively influenced by the solvent properties of water.
Hydrogen bonds and ionic, hydrophobic (Greek, "water-fearing"), and van der Waals interactions, although individually weak, are numerous in biological macromolecules and collectively have a very significant influence on the three-dimensional structures of proteins, nucleic acids, polysaccharides, and membrane lipids. Before we begin a detailed discussion of these biomolecules in the following chapters, it useful to review the properties of the solvent, water, in which they a assembled and carry out their functions.Weak Interactions in Aqueous Systems
| Hydrogen bonds between water molecules provide the cohesive forces that make water a liquid at room temperature and that favor the extreme ordering of molecules typical of crystalline water (ice). Polar biomolecules dissolve readily in water because they can replace energetically favorable water-water interactions with even more favorable water-solute interactions (hydrogen bonds and electrostatic interactions). In contrast, nonpolar biomolecules interfere with favorable water-water interactions and are poorly soluble in water. In aqueous solutions, these molecules tend to cluster together to minimize the energetically unfavorable effects of their presence. |
Hydrogen Bonding Gives Water Its Unusual Properties
Water has a higher melting point, boiling point, and heat of vapori~ tion than most other common liquids (Table 4-1). These unusual prc erties are a consequence of strong attractions between adjacent wat molecules, which give liquid water great internal cohesion.What is the cause of these strong intermolecular attractions liquid water? Each hydrogen atom of a water molecule shares an el~ tron pair with the oxygen atom. The geometry of the water molecule dictated by the shapes of the outer electron orbitals of the oxygen ato which are similar to the bonding orbitals of carbon (see Fig. 3-4 These orbitals describe a rough tetrahedron, with a hydrogen atom each of two corners and unshared electrons at the other two (Fig. 4-1) The H-O-H bond angle is 104.5° slightly less than the 109.5° of perfect tetrahedron; the nonbonding orbitals of the oxygen atc slightly compress the orbitals shared by hydrogen.
| The oxygen nucleus attracts electrons more strongly than does t hydrogen nucleus (i.e., the proton); oxygen is more electronegative (see Table 3-4). The sharing of electrons between H and O is therefc unequal; the electrons are more often in the vicinity of the oxygen atc than of the hydrogen. The result of this unequal electron sharing is two electric dipoles in the water molecule, one along each of the Hbonds; the oxygen atom bears a partial negative charge (δ-), and ea hydrogen a partial positive charge (δ+). The resulting electrostatic ; traction between the oxygen atom of one water molecule and the 1 drogen of another (Fig. 4-lc) constitutes a hydrogen bond. Hydrogen bonds are weaker than covalent bonds. The hydrog bonds in liquid water have a bond energy (the energy required break a bond) of only about 20 kJ/mol, compared with 460 kJ/mol i the covalent O-H bond. At room temperature, the thermal energy an aqueous solution (the kinetic energy resulting from the motion individual atoms and molecules) is of the same order as that requir to break hydrogen bonds. When water is heated, its temperate increase reflects the faster motion of individual water molecules. Although at any given time most of the molecules in liquid water are hydrogen-bonded, the lifetime of each hydrogen bond is less than 1 × 10~9 s. The apt phrase "flickering clusters" has been applied to the short-lived groups of hydrogen-bonded molecules in liquid water. The very large number of hydrogen bonds between molecules nevertheless confers great internal cohesion on liquid water. |
The nearly tetrahedral arrangement of the orbitals about the oxygen atom (Fig. 4-la) allows each water molecule to form hydrogen bonds with as many as four neighboring water molecules. At any given instant in liquid water at room temperature, each water molecule forms hydrogen bonds with an average of 3.4 other water molecules. The water molecules are in continuous motion in the liquid state, so hydrogen bonds are constantly and rapidly being broken and formed. In ice, however, each water molecule is fixed in space and forms hydrogen bonds with four other water molecules, to yield a regular lattice structure (Fig. 4-2). To break the large numbers of hydrogen bonds in such a lattice requires much thermal energy, which accounts for the relatively high melting point of water (Table 4-1). When ice melts orwater evaporates, heat is taken up by the system :
During melting or evaporation, the entropy of the aqueous system increases as more highly ordered arrays of water molecules relax into the less orderly hydrogen-bonded arrays in liquid water, or the wholly disordered water molecules in the gaseous state. At room temperature, both the melting of ice and the evaporation of water occur spontaneously; the tendency of the water molecules to associate through hydrogen bonds is outweighed by the energetic push toward randomness. ftecall that the free-energy change (ΔG) must have a negative value for a process to occur spontaneously: ΔG = ΔH - TΔS, where ΔG represents the driving force, ΔH the energy from making and breaking bonds, and ΔS the increase in randomness. Since ΔH is positive for melting and evaporation, it is clearly the increase in entropy (ΔS) that makes ΔG negative and drives these transformations.
Water Forms Hydrogen Bonds with Solutes
| Hydrogen bonds are not unique to water. They readily form between an electronegative atom (.usually oxygen or nitrogen) and a hydrogen atom covalently bonded to another electronegative atom in the same or another molecule (Fig. 4-3). However, hydrogen atoms covalently bonded to carbon atoms, which are not electronegative, do not participate in hydrogen bonding. The distinction explains why butanol (CH3CH2CH2CH2OH) has a relatively high boiling point of 117 'C, whereas butane (CH3CH2CH2CH3) has a boiling point of only -0.5 'C. Butanol has a polar hydroxyl group and thus can form hydrogen bonds with other butanol molecules. Uncharged but polar biomolecules such as sugars dissolve readily in water because of the stabilizing effect of the many hydrogen bonds that form between the hydroxyl groups or the carbonyl oxygen of the sugar and the polar water molecules. Alcohols, aldehydes, and ketones all form hydrogen bonds with water, as do compounds containing N-H bonds (Fig. 4-4), and molecules containing such groups tend to be soluble in water. |
Water Interacts Electrostatically with Charged Solutes
| Water is a polar solvent. It readily dissolves most biomolecules, whi are generally charged or polar compounds (Table 4-2); compounds th dissolve easily in water are hydrophilic (Greek, "water-loving"). contrast, nonpolar solvents such as chloroform and benzene are po solvents for polar biomolecules, but easily dissolve nonpolar biomo: cules such as lipids and waxes.
Water dissolves salts such as NaCl by hydrating and stabilizing the Na+ and Cl- ions, weakening their electrostatic interactions and thus counteracting their tendency to associate in a crystalline lattice (Fig. 4-6). The solubility of charged biomolecules in water is also a result of hydration and charge screening. Compounds with functional groups such as ionized carboxylic acids (-COO-), protonated aminines (-NH3 ), and phosphate esters or anhydrides are generally soluble in water for the same reason. Water is especially effective in screening the electrostatic interac- tions between dissolved ions. The strength, or force (F), of these ionic interactions depends upon the magnitude of the charges (Q), the dis- tance between the charged groups (r), and the dielectric constant (ε) of the solvent through which the interactions occur: |
The dielectric constant is a physical property reflecting the number dipoles in a solvent. For water at 25° C, ε (which is dimensionless) 78.5, and for the very nonpolar solvent benzene, e is 4.6. Thus ionic interactions are much stronger in less polar environments. The de- pendence on r2 is such that ionic attractions or repulsions operate over limited distances, in the range of 10 to 40 nm (depending on the elec- trolyte concentration) when the solvent is water.
Entropy Increases as Crystalline Substances Dissolve
As a salt such as NaCl dissolves, the Na+ and Cl- ions leaving the crystal lattice acquire far greater freedom of motion (Fig. 4-6). The resulting increase in the entropy (randomness) of the system is largely responsible for the ease of dissolving salts such as NaCI in water. In thermodynamic terms, formation of the solution occurs with a favorable change in free energy: ΔG = ΔH - TΔS, where ΔH has a small positive value and TΔS a large positive value; thus ΔG is negativeNonpolar Gases Are Poorly Soluble in Water
The biologically important gases C02, 02, and N2are nonpolar. In the diatomic molecules 02 and N2, electrons are shared equally by both atoms. In C02, each C=O bond is polar, but the two dipoles are oppositely directed and cancel each other (Table 4-3). The movement of these molecules from the disordered gas phase into aqueous solution constrains their motion and therefore represents a decrease in entropy. These gases are consequently very poorly soluble in water (Table 4-3). Some organisms have water-soluble carrier proteins (hemoglobin and myoglobin, for example) that facilitate the transport of O2. Carbon dioxide forms carbonic acid (H2C03) in aqueous solution, and is transported in that form.Two other gases, NH3 and H2S, also have biological roles in some organisms; these are polar and dissolve readily in water (Table 4-3).
