The chemical concept of pH measures the acidity or alkalinity of a solution on a scale typically ranging from 0 to 14. A stable pH environment is necessary for most chemical and biological processes to function correctly. Without this stability, reactions can slow down, speed up, or halt entirely. A buffer is a chemical system designed to prevent drastic fluctuations in pH when an acid or a base is introduced.
Defining the Chemical Components
A buffer solution is a specific mixture of two components that must be present in significant, balanced concentrations. This mixture consists of a weak acid and its corresponding conjugate base, or alternately, a weak base and its conjugate acid. This pairing is known as a conjugate acid-base pair.
The description “weak” means the acid or base does not fully dissociate into ions when dissolved in water, unlike strong acids or bases. For example, a weak acid, like acetic acid, remains largely intact in the solution, releasing only a small fraction of hydrogen ions ($\text{H}^{+}$). The conjugate base is what remains after the weak acid has released its hydrogen ion.
The Mechanism: How pH Change is Resisted
The buffer system resists pH change because its components are ready to absorb or neutralize any incoming acid or base. This action prevents the sharp rise or fall in hydrogen ion concentration that would otherwise occur. The two components of the buffer system have distinct roles in neutralizing these disturbances.
When a strong acid is added, it releases $\text{H}^{+}$ ions into the solution. The conjugate base component of the buffer immediately reacts with these added $\text{H}^{+}$ ions to form more of the original weak acid. Because the weak acid does not significantly dissociate, these newly formed molecules effectively capture the added acid, and the pH changes only minimally.
Conversely, if a strong base is added, it introduces hydroxide ions ($\text{OH}^{-}$) into the solution. The weak acid component of the buffer system neutralizes the base by donating its own $\text{H}^{+}$ ions to react with the added $\text{OH}^{-}$ ions, forming neutral water ($\text{H}_2\text{O}$). This reaction consumes the added base, keeping the $\text{pH}$ relatively stable.
Essential Roles in Living Systems
The maintenance of a stable $\text{pH}$ is important in biological systems, where chemical reactions are highly sensitive to their environment. Enzymes, which catalyze nearly all biological reactions, require a narrow $\text{pH}$ range to maintain their structure and function. A deviation outside this range can cause the enzymes to denature, rendering them ineffective and halting metabolic processes.
The bicarbonate buffer system is the primary mechanism that regulates the $\text{pH}$ of human blood and other extracellular fluids. This system consists of carbonic acid ($\text{H}_2\text{CO}_3$), the weak acid, and bicarbonate ion ($\text{HCO}_3^{-}$), its conjugate base. It works to maintain blood $\text{pH}$ within the narrow range of 7.35 to 7.45.
If the blood $\text{pH}$ drops below 7.35, acidosis occurs; a $\text{pH}$ above 7.45 leads to alkalosis. The bicarbonate buffer system is effective because its components are linked to the respiratory and renal systems, allowing the body to adjust acid and base levels rapidly. Excess acid is buffered by bicarbonate, and the resulting carbonic acid is converted into carbon dioxide ($\text{CO}_2$), which is expelled by the lungs. The body also uses the phosphate buffer system, which stabilizes the $\text{pH}$ inside cells (intracellular fluid).