Buffer solutions

When a weak acid and its conjugate base are both present in solution, small amounts of acid or base will result in a smaller pH change than would result if the same amount of acid or base was added to a neutral solution. Thus, a combination of weak acid and conjugate base is said to "buffer" the pH of a solution and so the solution is known as a buffer. For a solution of the weak acid $$HA$$ and its conjugate base $$A^-$$, addition of base or acid would result in one of the following:


 * $$HA+B\rightleftharpoons HB^+ +A^-$$
 * $$HA+OH^- \rightarrow H_2O+ A^-$$
 * $$A^-+H^+\rightleftharpoons HA$$

Using the example of an $$HA/A^-$$ solution, we have the following equilibrium-constant expression: $$

K_a=\frac{[A^-][H^+]}{[HA]}$$

Then, through algebraic manipulations, we have: $$\begin{array}{ccl} K_a&=&\frac{[A^-][H^+]}{[HA]}\\ \log K_a&=&\log \frac{[A^-][H^+]}{[HA]}\\ \log K_a&=&\log [A^-][H^+]-\log [HA]\\ \log K_a&=&\log [A^-]+\log [H^+] -\log [HA] \\ -\log [H^+]&=&-\log [K_a]+\log [A^-]-\log [HA]\\ pH&=&pK_a+\log \frac{[A^-]}{[HA]}\\ \end{array}$$

This equation is known as the Henderson-Hasselbalch equation, and it serves as a good way to approximate the pH of a buffer solution. If the $$K_a$$ of the acid is small enough, then the final concentrations of acid and base will be close to the same after equilibrium has been reached, and so using initial concentrations instead of final concentrations in the equation above produces a good approximation.