Living organisms rely on countless chemical reactions that are exquisitely sensitive to the concentration of hydrogen ions, measured as pH. Biological systems function optimally within an extremely narrow pH range. For example, human blood must be maintained between pH 7.35 and 7.45, a stability constantly challenged by metabolic activities that naturally produce acids. Buffers are the chemical defense system that resists these fluctuations, ensuring the stable internal environment necessary for life.
Why Precise pH Control is Essential for Life
The necessity for strict pH regulation stems from the structure and function of biological macromolecules, especially proteins and enzymes. Proteins are complex chains of amino acids that fold into precise three-dimensional shapes, and this shape dictates their function. The folding is stabilized by various chemical bonds, including ionic bonds between positively and negatively charged groups on the amino acid side chains.
Hydrogen ions from the surrounding fluid can interact with these charged groups, causing them to gain or lose a proton. A change in the overall charge profile of the protein disrupts the delicate ionic bonds that maintain its unique structure. This structural change, known as denaturation, prevents the protein from performing its specific job.
Enzymes, which catalyze reactions, are particularly susceptible to this effect. Each enzyme has an optimal pH at which its active site is correctly shaped to bind its molecule. The digestive enzyme pepsin, for example, functions best in the highly acidic environment of the stomach (pH 1.5–3.5), while enzymes in the small intestine prefer a more alkaline environment. Even a slight deviation from an enzyme’s optimal pH can reduce its efficiency or cause it to cease functioning entirely.
The Chemical Mechanism of Buffers
A buffer solution acts like a chemical shock absorber, minimizing the impact of added acid or base on the system’s pH. Chemically, a buffer is a mixture of a weak acid and its corresponding conjugate base, or a weak base and its conjugate acid. This pairing allows the solution to neutralize both added hydrogen ions (acids) and hydroxide ions (bases).
When an acid is introduced, the conjugate base component of the buffer system absorbs the excess hydrogen ions, converting them into a weak acid that does not significantly alter the pH. Conversely, if a base is added, the weak acid component of the buffer releases hydrogen ions to neutralize the incoming hydroxide ions, forming water. This dynamic equilibrium allows the system to absorb the stress of pH changes, keeping the hydrogen ion concentration relatively constant.
Major Buffer Systems in the Human Body
The human body employs several buffer systems to manage the pH in different fluid compartments. The most important of these is the bicarbonate buffer system, which is the primary mechanism for regulating the pH of the extracellular fluid, including the blood. This system relies on the reversible reaction between carbon dioxide (\(\text{CO}_2\)), water (\(\text{H}_2\text{O}\)), carbonic acid (\(\text{H}_2\text{CO}_3\)), and bicarbonate ions (\(\text{HCO}_3^-\)).
Carbon dioxide, a metabolic waste product, combines with water to form carbonic acid, which then dissociates into a hydrogen ion and a bicarbonate ion. This system is unique because its components are regulated by both the respiratory and renal systems.
The lungs can rapidly expel or retain \(\text{CO}_2\) to adjust the amount of acid in the blood within minutes. The kidneys provide a slower, but more powerful, long-term correction by excreting excess hydrogen ions into the urine. They also reabsorb or generate new bicarbonate ions to replenish the buffer supply.
The phosphate buffer system is important within the intracellular fluid and in the kidneys. It consists of dihydrogen phosphate (\(\text{H}_2\text{PO}_4^-\)), a weak acid, and hydrogen phosphate (\(\text{HPO}_4^{2-}\)), its conjugate base. This system is effective inside cells where phosphate concentrations are higher and plays a role in buffering urine. Proteins also function as buffers due to ionizable groups on their amino acid side chains, allowing them to bind or release hydrogen ions. Hemoglobin, a powerful protein buffer in red blood cells, binds to hydrogen ions produced from carbon dioxide metabolism as it releases oxygen, preventing a drop in blood \(\text{pH}\).
What Happens When Buffering Fails
When the body’s buffer capacity is overwhelmed, a condition of \(\text{pH}\) imbalance occurs, leading to a state of acidosis or alkalosis. Acidosis is defined as a blood \(\text{pH}\) that falls below 7.35, indicating an excess of hydrogen ions. Alkalosis is the opposite condition, where the blood \(\text{pH}\) rises above 7.45.
These conditions impair metabolic functions because the extreme \(\text{pH}\) levels cause protein and enzyme denaturation. For example, severe acidosis can depress the central nervous system, leading to confusion and coma. If the \(\text{pH}\) deviates too far from the physiological range, such as dropping below \(\text{pH}\) 6.8 or rising above \(\text{pH}\) 7.8, cellular processes are severely disrupted. This disruption can be rapidly fatal.