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From MyMCAT

Introduction

Buffer solutions are solutions capable of resisting change in pH. More specifically, they resist an increase in pH when Hydroxide is added, and they resist a decrease in pH when Hydronium is added. These solutions are ideal for reactions that must be carried out in precise pH conditions, especially when acids/bases are involved themselves, which is all too common in biological pathways.

A Thought Experiment to Understand Buffers

Suppose you had a solution of 1 part Sodium Acetate and 1 part Acetic acid. From our previous work we now understand that acetic acid, is in fact a weak acid which partially dissociates to form hydronium ions and acetate^- ions. We also know that this dissociation will be governed by the k_a of acetic acid forming a unique equilibrium at with specific concentrations. But now let us consider the Sodium Acetate, when dissolved, the sodium will do nothing (See previous section on Hydrolysis of Water due to salts) but the acetate^- ion on the other hand, has already been encountered. This excess of acetate ions will adjust, according to Le Chatelier's principle, in a shift backwards from the equilibrium.

How then does this solution, which we will consider to be only half-way dissociated towards products, react when either HCl or NaOH is added?

First let us consider adding HCl- a strong acid. HCl would increase the H⁺ concentration and, again according to Le Chatelier's principle, the equilibrium would shift away such that [H⁺] is lowered again. Thus what should have been a large decrease in pH (increase in acidity) was not that big because a lot of the H⁺ was consumed by the acetate ion.

If a strong base is added to a buffer, the weak acid will give up its H+ in order to transform the base (OH-) into water (H2O) and the conjugate base: HA + OH- → A- + H2O. Since the added OH- is consumed by this reaction, the pH will change only slightly.

Thus in either direction, the effects of strong acids and bases are reduced to only minimal changes in pH.

The Henderson-Hasselbalch Equation

From our previous work with acids and bases we know that we can easily write down a K_a expression for the equilibrium of the buffer.

The actual values, however, will have changed significantly from a simple dissociation because we have already added a portion of the conjugate base pair, but with some manipulation using logarithms we can arrive at the Henderson-Hasselbalch equation, which allows us to easily describe the pH in this new system.

In this equation,

• [A−] is the concentration of the conjugate base. This may be considered as coming completely from the salt, since the acid supplies relatively few anions compared to the salt lyn.
• [HA] is the concentration of the acid. This may be considered as coming completely from the acid, since the salt supplies relatively few complete acid molecules (A ⁻ may extract H ⁺ from water to become HA) compared to the added acid.

Maximum buffering capacity is found when pH = pKa, and the buffer's range is usually considered to be at a pH = pKa ± 1. Beyond this point, the ratio of acid to conjugate base fails to provide an effective buffering capability.