Читать книгу Geochemistry - William M. White - Страница 111

There is perhaps no compound more familiar to us than H₂O. Commonplace though it might be, H₂O is the most remarkable compound in nature. Its unusual properties include: the highest heat capacity of all solids and liquids except ammonia, the highest latent heat of vaporization of all substances, the highest surface tension of all liquids, its maximum density is at 4°C, with density decreasing below that temperature (negative coefficient of thermal expansion), the solid form is less dense than the liquid (negative Clapeyron slope), and finally, it is the best solvent known, dissolving more substances and in greater quantity than any other liquid. We will digress here briefly to consider the structure and properties of H₂O and the nature of water–electrolyte interactions from a microscopic perspective.

Many of the unusual properties of water arise from its nonlinear polar structure, which is illustrated in Figure 3.9a. The polar nature of water gives rise to van der Waals forces and the hydrogen bond discussed in Chapter 1. The hydrogen bond, which forms between hydrogen and oxygen atoms of adjacent molecules, imposes a dynamic partial structure on liquid water (Figure 3.9b). These bonds continually break and new ones reform, and there is always some fraction of unassociated molecules. On average, each water molecule is coordinated by four other water molecules. When water boils, all hydrogen bonds are broken. The energy involved in breaking these bonds accounts for the high heat of vaporization.

Figure 3.9 (a) Structure of the water molecule. Bond angle in the liquid phase is 108°, and 105° in the gas. The hydrogens retain a partial positive charge and the oxygen retains a partial positive charge. (b) Partial structure present in liquid water. Lines connecting adjacent molecules illustrate hydrogen bonds.

The dissolving power of water is due to its dielectric nature. A dielectric substance is one that reduces the forces acting between electric charges. When placed between two electrically charged plates (a capacitor), water molecules will align themselves in the direction of the electric field. As a result, the molecules oppose the charge on the plates and effectively reduce the transmission of the electric field. The permittivity, ε, of a substance is the measure of this effect. The relative permittivity, or dielectric constant, ε_r, of a substance is defined as the ratio of the capacitance observed when the substance is placed between the plates of a capacitor to the capacitance of the same capacitor when a vacuum is present between the plates:

(3.67)

where ε₀ is the permittivity of a vacuum (8.85 × 10⁻¹² C²/J m). The relative permittivity of water is 78.54 at 25°C and 1 atm. For comparison, the relative permittivity of methane, a typical nonpolar molecule, is 1.7.

Water molecules surrounding a dissolved ion will tend to align themselves to oppose the charge of the ion. This insulates the ion from the electric field of other ions. This property of water accounts in large measure for its dissolving power. For example, we could easily calculate that the energy required to dissociate NaCl (i.e., the energy required to move Na⁺ and Cl^– ions from their normal interatomic distance in a lattice, 236 pm, to infinite separation) is about 585 kJ/mol. Because water has a dielectric constant of about 80, this energy is reduced by a factor of 80, so only 7.45 kJ are required for dissociation.

The charged nature of ions and the polar nature of water result in the solvation of dissolved ions. Immediately adjacent to the ion, water molecules align themselves to oppose the charge on the ion, such that the oxygen of the water molecule will be closest to a cation (Figure 3.10). These water molecules are called the first solvation shell or layer and they are effectively bound to the ion, moving with it as it moves. Beyond the first solvation shell is a region of more loosely bound molecules that are only partially oriented, called the second solvation shell or layer. The boundary of this latter shell is diffuse: there is no sharp transition between oriented and unaffected water molecules. The energy liberated in this process, called the solvation energy, is considerable. For NaCl, for example, it is −765kJ/mol (it is not possible to deduce the solvation energies of Na⁺ and Cl^– independently). The total number of water molecules bound to the ion is called the solvation number. Solvation effectively increases the electrostatic radius of cations by about 90 pm and of anions by about 10 pm per unit of charge.

Figure 3.10 Solvation of a cation in aqueous solution. In the first solvation shell, water molecules are bound to the cation and oriented so that the partial negative charge on the oxygen faces the cation. In the second solvation shell, molecules are only loosely bound and partially oriented.

An additional effect of solvation is electrostriction. Water molecules in the first solvation sphere are packed more tightly than they would otherwise be. This is true, to a lesser extent, of molecules in the secondary shell. In addition, removal of molecules from the liquid water structure causes partial collapse of this structure. The net effect is that the volume occupied by water in an electrolyte solution is less than in pure water, which can lead to negative apparent molar volumes of solutes, as we shall see. The extent of electrostriction depends strongly on temperature and pressure.

A final interesting property of water is that some fraction of water molecules will autodissociate. In pure water at standard state conditions, one in every 10⁻⁷ molecules will dissociate to form H⁺ and OH^– ions. Although in most thermodynamic treatments the protons produced in this process are assumed to be free ions, most will combine with water molecules to form H₃O⁺ ions. OH⁻ is called the hydroxyl ion; the H₃O⁺ is called hydronium.