A measure of acidity or basicity, called pH, is a fundamental property of any aqueous solution, determined by the concentration of hydrogen ions (\(H^+\)). The pH scale ranges from 0 to 14: values below 7 are acidic, 7 is neutral, and values above 7 are basic. Maintaining a stable pH is important for biological systems, such as the human body, because slight fluctuations can disrupt the structure and function of proteins and enzymes. For instance, human blood must remain within the narrow range of 7.35 to 7.45. A chemical system called a buffer resists significant changes in pH when small amounts of acid or base are added by neutralizing incoming substances to keep the hydrogen ion concentration relatively constant.
The Essential Components of a Buffer
A buffer solution consists of a conjugate pair: a weak acid and its corresponding conjugate base, or a weak base and its conjugate acid. Both components must be present in appreciable concentrations for the buffer to function effectively. A weak acid, such as acetic acid, only partially dissociates in water. Its conjugate base, the acetate ion, is the remaining chemical structure once the weak acid has donated a proton.
This two-part chemical team is the core of the buffer system, allowing it to respond to both increasing acidity and increasing basicity. The weak acid component neutralizes any added base, while the conjugate base component neutralizes any added acid. The specific ratio of these two components determines the buffer’s resting pH.
How Buffers Handle Added Acid
When a strong acid is introduced, it releases excess hydrogen ions (\(H^+\)). The conjugate base component of the buffer system immediately reacts with these free hydrogen ions. This reaction converts the strong acid into a molecule of the buffer’s weak acid component.
For example, in the acetic acid/acetate buffer, the acetate ion (\(CH_3COO^-\)) reacts with the added \(H^+\) to form more acetic acid (\(CH_3COOH\)). Since acetic acid is a weak acid, it remains largely undissociated. By transforming the highly acidic hydrogen ions into a much weaker acid, the buffer prevents a significant change in the overall concentration of free \(H^+\) in the solution.
How Buffers Handle Added Base
When a strong base, which releases hydroxide ions (\(OH^-\)), is added, the weak acid component of the buffer system counteracts the change. The weak acid readily donates a proton (\(H^+\)) to the incoming hydroxide ion.
This reaction neutralizes the strong base, combining the \(H^+\) from the weak acid with the \(OH^-\) from the added base to form a neutral water molecule (\(H_2O\)). The result of this neutralization is the formation of the buffer’s conjugate base, slightly increasing its concentration. By converting the highly reactive hydroxide ions into water, the weak acid component prevents a drastic jump in the solution’s pH.
Buffer Capacity and Biological Importance
The ability of a buffer to absorb added acid or base is limited; this limit is known as its buffer capacity. Buffer capacity is defined as the amount of strong acid or strong base that can be added before the pH begins to change significantly. Once one of the buffer components—either the weak acid or the conjugate base—is nearly exhausted, the system becomes overwhelmed, and the pH will change rapidly. Capacity is highest when the concentrations of the weak acid and its conjugate base are approximately equal.
In living organisms, buffer systems maintain homeostasis, the stable internal environment necessary for life. The bicarbonate buffer system is a primary example, playing a major role in regulating the pH of human blood. This system involves carbonic acid (\(H_2CO_3\)), the weak acid, and bicarbonate ions (\(HCO_3^-\)), the conjugate base.
It is effective because carbonic acid is in equilibrium with carbon dioxide (\(CO_2\)), a gas that can be quickly expelled or retained by the lungs. This interconnectedness with the respiratory system allows the body to continuously replenish the buffer components, giving the bicarbonate system a high practical capacity in vivo. When metabolism produces excess acid, the bicarbonate ions neutralize it, forming carbonic acid, which is then converted to \(CO_2\) and exhaled. This rapid removal of the reaction product helps keep the blood pH tightly regulated, preventing life-threatening conditions like acidosis or alkalosis.