The acidity or alkalinity of a solution is measured by the pH scale, which quantifies the concentration of hydrogen ions (\(\text{H}^+\)). Even a small change in the concentration of these ions can cause a dramatic shift in the chemical environment. A pH buffer is a chemical solution designed to resist significant changes to this acidity or alkalinity when a small amount of strong acid or strong base is introduced. This resistance is achieved by neutralizing the added acid or base, stabilizing the system.
What a Buffer System is Made Of
A functional pH buffer is not a single chemical substance but a mixture of two specific components: a weak acid and its corresponding conjugate base, or a weak base and its corresponding conjugate acid. A weak acid, such as acetic acid, does not fully dissociate in water, meaning most molecules remain intact. The conjugate base is the substance formed when the weak acid loses a hydrogen ion.
For example, a common laboratory buffer pairs acetic acid with its conjugate base, the acetate ion, often supplied as a salt like sodium acetate. The buffer requires a reservoir of both the weak acid and its conjugate base in equilibrium to neutralize both incoming acids and bases.
How Buffers Resist pH Change
The mechanism of a buffer’s action is a two-sided neutralization process that maintains the existing chemical balance.
Neutralizing Added Acid
When a strong acid, which releases hydrogen ions (\(\text{H}^+\)), is added, the conjugate base component immediately reacts with these ions. This reaction “mops up” the newly added \(\text{H}^+\) by forming more of the original weak acid. Since the weak acid does not significantly contribute to the overall acidity, the \(\text{pH}\) remains stable.
Neutralizing Added Base
Conversely, if a strong base is added, it introduces hydroxide ions (\(\text{OH}^-\)) that would ordinarily raise the \(\text{pH}\) dramatically. The weak acid component of the buffer system donates its own hydrogen ions to the added hydroxide ions. This acid-base reaction neutralizes the strong base, resulting in the formation of water and more of the conjugate base.
The buffer components convert the highly reactive strong acid or strong base into a much weaker substance. This chemical conversion minimizes the shift in the concentration of free hydrogen ions, which determines the solution’s \(\text{pH}\). The capacity of a buffer—how much acid or base it can neutralize before the \(\text{pH}\) changes significantly—is determined by the concentration of the weak acid and conjugate base present.
Buffers in Living Systems
Buffer systems are necessary for maintaining the internal environment required for life, a condition known as homeostasis. In human physiology, the most prominent example is the bicarbonate buffer system, which keeps blood \(\text{pH}\) within a tightly controlled range of 7.35 to 7.45. A blood \(\text{pH}\) outside this narrow window, such as below 6.8 or above 7.8, can be lethal because it interferes with the structure and function of proteins and enzymes.
The bicarbonate buffer system consists of carbonic acid (\(\text{H}_2\text{CO}_3\)) and its conjugate base, bicarbonate (\(\text{HCO}_3^-\)). When metabolic processes generate excess acid, such as lactic acid, the bicarbonate ions neutralize the added \(\text{H}^+\) to form carbonic acid. This carbonic acid then dissociates into water and carbon dioxide (\(\text{CO}_2\)), which the lungs rapidly expel through increased breathing.
If the blood becomes too alkaline, the carbonic acid component neutralizes the excess base, shifting the system back toward a lower \(\text{pH}\). This equilibrium is important because enzymes have optimal activity levels within a specific \(\text{pH}\) range. Failure leads to conditions like acidosis or alkalosis, which severely impair bodily functions.
Everyday Uses for pH Buffers
Buffer solutions extend their stabilizing function into numerous industrial and commercial applications.
- In the pharmaceutical industry, buffers maintain the precise \(\text{pH}\) necessary for drug stability and effectiveness, often extending shelf life.
- The food and beverage industry uses buffers to control acidity, which affects flavor, texture, and preservation. Common additives help maintain a consistent \(\text{pH}\) in dairy products, canned goods, and during fermentation.
- In cosmetics and personal hygiene, products are buffered to maintain a \(\text{pH}\) compatible with the skin and scalp (typically around \(\text{pH}\) 5.5) to prevent irritation.
- Buffers are used in laboratory settings for the calibration of \(\text{pH}\) meters and to create controlled environments for chemical experiments.