The acidity or alkalinity of a solution is measured using the pH scale, which quantifies the concentration of hydrogen ions (\(H^+\)) present. A low pH indicates a high concentration of these ions, making the solution acidic, while a high pH indicates a low concentration, resulting in an alkaline or basic solution. In many chemical reactions and especially within biological systems, a stable pH environment is necessary for processes to function correctly. A buffer solution is a chemical system specifically designed to resist drastic shifts in this hydrogen ion concentration when an acid or a base is introduced. This resistance to change allows delicate systems, such as human blood, to maintain a tightly controlled pH range (7.35 to 7.45), preventing metabolic disruption.
What Buffer Solutions Are Made Of
A functional buffer system is composed of a pair of related chemical compounds: a weak acid and its corresponding conjugate base, or a weak base and its corresponding conjugate acid. This pairing is crucial because it provides both a component to react with added acid and a component to react with added base. The term “weak” is important because it means the acid or base only partially dissociates into ions when dissolved in water, maintaining a significant reservoir of the original, unreacted molecule.
For example, an acidic buffer might be created by combining acetic acid, a weak acid, with sodium acetate, which provides the acetate ion, its conjugate base. The weak acid component is ready to donate a proton, while the conjugate base is prepared to accept one. Conversely, a basic buffer would involve a weak base, like ammonia, and its conjugate acid, ammonium chloride. In either case, the two components exist in a state of chemical equilibrium, ready to counter any external disturbance to the pH.
The Chemical Mechanism of pH Stabilization
The ability of a buffer to stabilize pH is rooted in two separate reactions that neutralize any added acid (\(H^+\)) or base (\(OH^-\)). When hydrogen ions are added, the buffer’s conjugate base component acts as a sponge to absorb them. Consider the acetic acid/acetate buffer: the acetate ion (\(CH_3COO^-\)) reacts with the added \(H^+\) ions to form more of the weak acid, acetic acid (\(CH_3COOH\)).
\(CH_3COO^- \text{ (Conjugate Base)} + H^+ \text{ (Added Acid)} \rightarrow CH_3COOH \text{ (Weak Acid)}\)Because acetic acid is a weak acid, it remains mostly undissociated and contributes very few new hydrogen ions to the solution. This effectively removes the highly reactive, added \(H^+\) from the solution, preventing a large drop in pH.
When a base is added, the buffer’s weak acid component takes over the neutralizing role. The added hydroxide ions (\(OH^-\)) from the base are highly reactive and would normally cause a sharp increase in pH. The weak acid, such as acetic acid, readily donates its proton to the hydroxide ion.
\(CH_3COOH \text{ (Weak Acid)} + OH^- \text{ (Added Base)} \rightarrow H_2O \text{ (Water)} + CH_3COO^- \text{ (Conjugate Base)}\)This reaction converts the strong base (\(OH^-\)) into neutral water and the buffer’s own conjugate base, which is a much weaker base. By neutralizing the hydroxide ions, the weak acid component prevents the rapid increase in alkalinity.
Limits of Buffer Action
A buffer solution has limitations described by buffer capacity and buffer range. Buffer capacity refers to the maximum quantity of strong acid or strong base that can be added to the solution before a significant change in pH is observed.
The capacity is directly related to the concentration of the buffer components. A solution with a higher concentration of the weak acid and its conjugate base will have a greater capacity to neutralize additions than a dilute solution. Once the quantity of added acid or base exceeds the available amount of the neutralizing buffer component, the buffering action stops, and the pH begins to change rapidly.
Buffer range defines the specific pH interval over which the buffer system can effectively resist changes. A buffer works most efficiently when the concentration of the weak acid is approximately equal to the concentration of its conjugate base. This optimal zone of activity generally spans a range of one pH unit above and one pH unit below the acid’s pKa value. The pKa is a constant specific to the weak acid; selecting a buffer with a pKa close to the desired operational pH ensures the system is effective.