Much important chemistry, including almost all of the chemistry of the natural world, occur in aqueous solution. We have studied one very significant class of aqueous equilibria, acid-base reactions. Now we will consider applications of acid-base chemistry and study two additional types of aqueous equilibria, those involving the solubility of salts and those involving the formation of complex ions. The interaction of acid-base, solubility, and complex ion equilibria is often important in natural processes, such as the weathering of minerals, the uptake of nutrients by plants, and tooth decay.

Solutions of Acids or Bases Containing a Common Ion Consider a solution containing the weak acid hydrofluoric acid (HF, K a = 7.2 x ) and its salt sodium fluoride (NaF). Remember when a salt dissolves in water, it breaks up into its ions (it is a strong electrolyte). NaF(s) → Na + (aq) + F - (aq)(in water) HF is a weak acid so the major species are HF, Na +, F -, and H 2 O. The common ion in this solution is F - since it is produced by HF and NaF.

What effect does F - from NaF have on the dissociation equilibrium of HF? By Le Chatelier’s principle, the following occurs: HF(aq) ⇌ H + (aq) + F - (aq) Equilibrium shifts away from added component. Fewer H + ions present. Added F - ions from NaF The shift in equilibrium that occurs because of the addition of an ion already involved in the equilibrium is called the common ion effect. Makes the NaF and HF solution less acidic than HF solution alone.

The most important application of acid-base solutions containing a common ion is for buffering. A buffered solution is one that resists a change in its pH when either hydroxide or protons are added. A buffered solution may contain a weak acid and its salt (for example, HF and NaF) or a weak base and its salt (for example, NH 3 and NH 4 Cl).

The most important practical example of a buffered solution is our blood, which can absorb the acids and bases produced in biological reactions without changing its pH. A constant pH for blood is vital because cells can survive only in a very narrow pH range (7.35 to 7.45).

Suppose a buffered solution contains relatively large amounts of a weak acid HA and its conjugate base A -. When a strong base (OH - ) is added to the solution, since the weak acid is the best source of protons (H + ), the following reaction occurs: OH - + HA → A - + H 2 O The net result is that OH - ions are not allowed to accumulate but are replaced by A - ions. Original buffer pH OH - added Final pH of buffer close to original Added OH - ions replaced by A - ions

To understand the stability of the pH under these conditions examine the equilibrium expression for HA: rearranging gives Since pH is governed by [H + ], in a buffered solution the pH is dependent on the ratio of [HA]/[A - ]. When OH - ions are added, HA is converted to A -, and the ratio [HA]/[A - ] decreases. But if the amount of HA and A - are very large compared with the OH - added, the change in the [HA]/[A - ] ratio will be small and the [H + ] and pH remain essentially unchanged.

If protons (H + ) are added to a buffered solution of a weak acid and a salt of its conjugate base, the added protons react with A - to form weak acid, HA H + + A - → HA and free H + ions do not accumulate. If [A - ] and [HA] are large compared with the [H + ] added, little change in the pH will occur.

The buffering capacity of a buffered solution represents the amount of H + or OH - ions the buffer can absorb without a significant change in pH. A buffer with a large capacity contains large concentrations of buffering components and so can absorb a relatively large amount of H + or OH - ions and show little pH change. The capacity of a buffered solution is determined by the magnitudes of [HA] and [A - ].

The mole quantities of buffer components have to be considerable greater – upwards to 10 times greater – than the expected influx of strong acid or base if the change in the pH of the system is to be kept with 0.1 pH units of the original.

Optimal buffering occurs when [HA] is equal to [A - ]. This means that when choosing a buffer for a specific application, we want [A - ]/[HA] to equal 1. Since the pK a of the weak acid to be used in the specific buffer should be as close as possible to the desired pH ( pH = pK a ± 1).