Question

A buffer system used to maintain solution pH values around 5 is prepared by adding similar concentrations of \(\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}\) and \(\mathrm{NaCH}_{3} \mathrm{CO}_{2}\) to water. (a) Write a balanced equation showing the acid and conjugate base in equilibrium. Omit spectator ions. (b) Describe how this buffer system prevents large \(\mathrm{pH}\) changes when a base is added.

Short answer

The balanced equation for the equilibrium of the weak acid (acetic acid) and its conjugate base (acetate ion) is \(\mathrm{CH}_{3}\mathrm{CO}_{2}\mathrm{H} \leftrightarrow \mathrm{H}^{+} + \mathrm{CH}_{3}\mathrm{CO}_{2}^{-}\). When a base is added to the buffer system, even though it tends to increase the pH by removing \(\mathrm{H}^{+}\) ions, the acetic acid in the buffer system replaces the \(\mathrm{H}^{+}\) ions, thereby keeping the \(\mathrm{pH}\) almost constant.

Step-by-step solution

Write a balanced equation

The buffer system consists of a weak acid, acetic acid \(\mathrm{CH}_{3}\mathrm{CO}_{2}\mathrm{H}\), and its conjugate base, the acetate ion \(\mathrm{CH}_{3} \mathrm{CO}_{2}^{-}\), which comes from its salt, sodium acetate \(\mathrm{NaCH}_{3} \mathrm{CO}_{2}\). The balanced equation, omitting the spectator ions, can be written as following: \[\mathrm{CH}_{3}\mathrm{CO}_{2}\mathrm{H} \leftrightarrow \mathrm{H}^{+} + \mathrm{CH}_{3}\mathrm{CO}_{2}^{-}\]

Describe the buffering action when a base is added

When a base is added to the buffer system, it tends to increase the pH by removing \(\mathrm{H}^{+}\) ions. However, the acetic acid present in the buffer system acts to counteract this. It will donate more of its own \(\mathrm{H}^{+}\) ions to replace those removed by the base. As a result, the concentration of \(\mathrm{H}^{+}\) ions remains almost constant and the \(\mathrm{pH}\) changes very little. Thus, the buffer system prevents large shifts in \(\mathrm{pH}\) when a base is added.

Key concepts

Acid-Base Equilibrium

Understanding acid-base equilibrium is key to grasping why buffer systems are so important in chemistry. At the heart of this concept is the reversible reaction wherein an acid donates protons (hydrogen ions, \( \mathrm{H}^{+} \) to the solution, and the base accepts them.

For instance, when acetic acid (\(\mathrm{CH}_{3}\mathrm{CO}_{2}\mathrm{H}\)) is dissolved in water, it partially dissociates into hydrogen ions and acetate ions (\(\mathrm{CH}_{3}\mathrm{CO}_{2}^{-}\)). This reaction reaches a state of equilibrium, where the rate of the forward reaction (the acid dissociating) is equal to the rate of the reverse reaction (the acid reforming from its ions). This balance is critical to a buffer's ability to maintain a stable pH, as it provides a reservoir of ions that can react with any added acid or base.

\[\mathrm{CH}_{3}\mathrm{CO}_{2}\mathrm{H} \leftrightarrow \mathrm{H}^{+} + \mathrm{CH}_{3}\mathrm{CO}_{2}^{-}\]

Buffer Action

Buffer action is the mechanism by which buffers resist changes in pH. A buffer typically consists of a weak acid and its conjugate base, or a weak base and its conjugate acid. This allows the buffer to neutralize added acids or bases.

In the case of the acetic acid and acetate buffer, when a strong base is added to the solution, the acetate ions (\(\mathrm{CH}_{3}\mathrm{CO}_{2}^{-}\)) react with the added base's \(\mathrm{H}^{+}\) ions. Acetic acid can donate its \(\mathrm{H}^{+}\) ions to any residual base, reforming water and ensuring that the pH remains constant. If an acid is added instead, the acetate ions will react with the \(\mathrm{H}^{+}\) from the acid, forming more acetic acid and mitigating the pH change.

This elegant chemical dance is why buffers are the unsung heroes in many biological and chemical systems, maintaining conditions necessary for processes to occur within a narrow and optimal pH range.

pH Regulation

The regulation of pH is a fundamental aspect of chemistry, especially in biological systems where enzymes and other macromolecules require specific pH levels to function properly. pH is defined as the negative logarithm of the \(\mathrm{H}^{+}\) ion concentration, and small changes in \(\mathrm{H}^{+}\) levels can lead to significant shifts in pH.

Buffers play an essential role in the regulation of pH. They work by employing a combination of a weak acid or base and its salt, which consist of a conjugate base or acid. This combination allows the solution to withstand drastic pH changes, even when strong acids or bases are introduced. The strength of a buffer is determined by the concentrations of the weak acid and its conjugate base and by the acid's dissociation constant (\(Ka\)). The closer the pH is to the \(\mathrm{p}\Ka\) of the acid, the more effective the buffer will be in resisting pH changes.

Through the thoughtful application of buffer systems, chemists can create environments that simulate the tightly regulated pH conditions found within living organisms, opening the door to myriad research and industrial applications.