### Introduction

A buffer is an aqueous mixture of a weak acid and its conjugate base, or a weak base and its conjugate acid. A buffer has a very stable pH. When small amounts of an acid or a base are added to a buffer solution, the pH of the solution changes very little. In many chemical and biochemical systems, buffers are critical. Blood plasma, a natural example in humans, is a bicarbonate buffer that keeps the pH of blood between 7.2 and 7.6.

By design, a buffer is an equilibrium system. An example is a buffer prepared with nitrous acid, HNO2. The weak acid establishes an aqueous equilibrium as shown below: ${\text{HNO}}{_{\text{2}}}{\text{ (aq)}} \to {\text{H}}{^{+}}{\text{ (aq) + NO}}{_{\text{2}}}{^{\text{-}}}{\text{ (aq)}}$

The equilibrium constant expression is $K{_{a}} = \frac{[{\text{H}}{^{+}}][{\text{NO}}{_{2}}{^{-}}]}{[{\text{HNO}}{_{2}}]}$

To prepare a buffer system with nitrous acid, a conjugate base, such as sodium nitrite (NaNO2), is added. The resulting system is a mixture of HNO2 and NO2 ions. The nitrous acid molecule will buffer the addition of an acid and the nitrite ion from the conjugate will buffer the addition of a base.

A variation of the equilibrium expression above, called the Henderson-Hasselbalch equation, is a very good reference in preparing a buffer solution. For a nitrous acid/sodium nitrate buffer, the Henderson-Hasselbalch equation is shown below: $pH = pK{_{a}} + log \frac{[{\text{NO}}{_{2}}{^{-}}]}{[{\text{HNO}}{_{2}}]}$

The pH range in which a buffer solution is effective is generally considered to be ±1 of the pKa.

### Objectives

In the Initial Investigation, you will prepare two acetic acid/sodium acetate buffer systems. After examining the factors involved in preparing these buffers and testing their effectiveness, you will design a buffer for a specified pH range as well as test its effectiveness; effectiveness will be assessed based on the amount of acid or base a buffer can absorb before the pH of the buffer solution falls outside a specified pH range.