0.21 Acid ionization equilibrium  (Page 6/7)

Unlike our previous NaCl salt solution, a measurement in this case reveals that the pH of the product salt solution is 9.4, indicating that the solution is basic. Thus, mixing equal molar quantities of strong base with weak acid produces a basic solution. In essence, the weak acid does not fully neutralize the strong base. To understand this, we examine the behavior of sodium acetate in solution. Since the pH is greater than 7, then there is an excess of OH ^- ions in solution relative to pure water. These ions must have come from the reaction of sodium acetate with the water. Therefore, the negative acetate ions in solution must behave as a base, accepting positive hydrogen ions:

A ^- (aq) + H ₂ O(l) → HA(aq) + OH ^- (aq)

The reaction of an ion with water to form either an acid or a base solution is referred to as "hydrolysis." From this example, the salt of a weak acid behaves as a base in water, resulting in a pH greater than 7.

To understand the extent to which the hydrolysis of the negative ion occurs, we need to know the equilibrium constant for this reaction. This turns out to be determined by the acid ionization constant for HA. To see this, we write the equilibrium constant for the hydrolysis of A ^- as

${\text{K}}_{\text{h}}=\frac{\left[\text{H}\text{A}\right]\left[\text{O}{\text{H}}^{\text{-}}\right]}{\left[{\text{A}}^{\text{-}}\right]}$

Multiplying numerator and denominator by [H ₃ O ⁺ ], we find that

${\text{K}}_{\text{h}}=\frac{\left[\text{H}\text{A}\right]\left[\text{O}{\text{H}}^{\text{-}}\right]}{\left[{\text{A}}^{\text{-}}\right]}\frac{\left[{\text{H}}_{\text{3}}{\text{O}}^{+}\right]}{\left[{\text{H}}_{\text{3}}{\text{O}}^{+}\right]}=\frac{\left[{\text{K}}_{\text{w}}\right]}{\left[{\text{K}}_{\text{a}}\right]}$

Therefore, for the hydrolysis of acetate ions in solution, K _h = 5.8∙10 ^-10 . This is fairly small, so the acetate ion is a very weak base.

Observation 5: acid strength and molecular properties

We now have a fairly complete quantitative description of acid-base equilibrium. To complete our understanding of acid-base equilibrium, we need a predictive model which relates acid strength or base strength to molecular properties. In general, we expect that the strength of an acid is related either to the relative ease by which it can donate a hydrogen ion or by the relative stability of the remaining negative ion formed after the departure of the hydrogen ion.

To begin, we note that there are three basic categories of acids that we have examined in this study. First, there are simple binary acids: HF, HCl, HBr, and HI. Second, there are acids formed from main group elements combined with one or more oxygen atoms, such H ₂ SO ₄ or HNO ₃ . These are called "oxyacids." Third, there are the "carboxylic acids," organic molecules which contain the carboxylic functional group:

We consider first the simple binary acids. HCl, HBr, and HI are all strong acids, whereas HF is a weak acid. In comparing the experimental values of pK _a values in Table 7, we note that the acid strength increases in the order HF<HCl<HBr<HI. This means that the hydrogen ion can more readily separate from the covalent bond with the halogen atom (X) as we move down the periodic table. This is reasonable, because the strength of the H-X bond also decreases as we move down the periodic table, as shown in Table 7:

The decreasing strength of the H-X bond is primarily due to the increase is the size of the X atom as we move down the periodic table. We conclude that one factor which influences acidity is the strength of the H-X bond: a weaker bond produces a stronger acid, and vice versa.