Question

Briefly describe how a buffer solution can control the pH of a solution when strong acid is added and when strong base is added. Use \(\mathrm{NH}_{3} / \mathrm{NH}_{4} \mathrm{Cl}\) as an example of a buffer and \(\mathrm{HCl}\) and \(\mathrm{NaOH}\) as the strong acid and strong base.

Short answer

The buffer absorbs \(\mathrm{H}^+\) from acids to \(\mathrm{NH}_4^+\) and \(\mathrm{OH}^-\) from bases to \(\mathrm{NH}_3\) , stabilizing pH.

Step-by-step solution

Define Buffer Solution

A buffer solution consists of a weak acid and its conjugate base or a weak base and its conjugate acid. In this case, the buffer is composed of ammonia (\(\mathrm{NH}_3\)) and ammonium chloride (\(\mathrm{NH}_4\mathrm{Cl}\)). It resists changes in pH upon the addition of small amounts of strong acids or bases.

Addition of Strong Acid

When a strong acid such as hydrogen chloride (\(\mathrm{HCl}\)) is added to the\(\mathrm{NH}_3 / \mathrm{NH}_4\mathrm{Cl}\) buffer solution, the hydrogen ions (\(\mathrm{H}^+\)) from the\(\mathrm{HCl}\) are neutralized by the ammonia, which acts as a weak base. The ammonia reacts to form ammonium ions (\(\mathrm{NH}_4^+\)), thus reducing the effect of the added \(\mathrm{H}^+\) ions: \[ \mathrm{NH}_3 + \mathrm{H}^+ \rightarrow \mathrm{NH}_4^+ \] This reaction minimizes the increase in hydrogen ion concentration, hence resisting a decrease in pH.

Addition of Strong Base

When a strong base like sodium hydroxide (\(\mathrm{NaOH}\)) is added to the buffer, the hydroxide ions (\(\mathrm{OH}^-\)) react with the ammonium ions (\(\mathrm{NH}_4^+\)) present in the buffer to form ammonia and water: \[ \mathrm{NH}_4^+ + \mathrm{OH}^- \rightarrow \mathrm{NH}_3 + \mathrm{H}_2\mathrm{O} \] This reaction reduces the \(\mathrm{OH}^-\) concentration, minimizing the increase in pH that would typically occur with the addition of a base.

Buffer Mechanism Summary

The \(\mathrm{NH}_3 / \mathrm{NH}_4Cl\) buffer effectively controls pH by converting \(\mathrm{H}^+\) from strong acids into \(\mathrm{NH}_4^+\) ions, and \(\mathrm{OH}^-\) from strong bases into \(\mathrm{NH}_3\) and water. It maintains equilibrium by using the weak base \(\mathrm{NH}_3\) and its conjugate acid \(\mathrm{NH}_4^+\) to absorb excess \(\mathrm{H}^+\) or \(\mathrm{OH}^-\) ions, hence stabilizing the solution's pH.

Key concepts

pH control

Buffer solutions are a fascinating aspect of chemistry that help keep the pH level of a solution stable, even when acids or bases are added. They are integral in various applications such as maintaining the right environment in biological systems or chemical reactions. A buffer works by using either a weak acid and its conjugate base or a weak base and its conjugate acid. This unique combination absorbs excess hydrogen ions (\(\mathrm{H}^+\)) or hydroxide ions (\(\mathrm{OH}^-\)), which are introduced when strong acids or bases are added to the solution. This absorption prevents drastic changes in pH, effectively 'buffering' the solution from becoming too acidic or basic.

pH control is crucial in many areas like biological processes, where precise pH levels are necessary for proper function. For example, in the human body, blood pH is tightly regulated by buffer systems to stay around 7.4. Such control ensures enzymes and cellular processes operate optimally. The ability of a buffer to maintain pH cannot be overstated, as it allows for continuity in processes that are dependent on specific pH conditions.

acid-base reactions

Acid-base reactions form the core of many chemical processes and the principle behind buffer systems. When acids and bases interact, they can neutralize each other. For instance, when an acid such as hydrochloric acid (\(\mathrm{HCl}\)) is added to a buffer, it releases hydrogen ions (\(\mathrm{H}^+\)). These ions are then "captured" by the weak base present in the buffer, such as ammonia (\(\mathrm{NH}_3\)), forming ammonium ions (\(\mathrm{NH}_4^+\)). This reaction reduces the concentration of free hydrogen ions, preventing a significant decline in pH.

Similarly, when a base like sodium hydroxide (\(\mathrm{NaOH}\)) is added, it provides hydroxide ions (\(\mathrm{OH}^-\)). These ions react with the conjugate acid in the buffer, such as ammonium (\(\mathrm{NH}_4^+\)), to produce water and regenerate ammonia. This mechanism decreases the free hydroxide ion concentration, thus avoiding a rise in pH. Hence, acid-base reactions in buffer systems are not just about neutralization, but about stabilizing a solution's pH to remain constant amidst changes.

ammonia-ammonium buffer

The ammonia-ammonium buffer is a prime example of a pH-regulating system utilizing a weak base and its conjugate acid. It is specifically composed of ammonia (\(\mathrm{NH}_3\)), a weak base, and ammonium chloride (\(\mathrm{NH}_4\mathrm{Cl}\)), which supplies ammonium ions (\(\mathrm{NH}_4^+\)). When these two components are in solution together, they maintain the pH by reacting with any added acid or base.

This buffer is particularly effective because the reactions that consume either the excess hydrogen ions from strong acids or the excess hydroxide ions from strong bases do not go to completion. Instead, they shift the equilibrium, allowing the solution to resist changes in pH. For example, upon addition of \(\mathrm{HCl}\), the ammonia reacts with the hydrogen ions, converting them to ammonium ions. Conversely, if \(\mathrm{NaOH}\) is introduced, the ammonium ions react with the hydroxide ions, reforming ammonia and creating water. Thus, the ammonia-ammonium buffer demonstrates how delicate chemical balance can provide significant control over a solution's pH in dynamic conditions.