The measure of a solution’s acidity or basicity, known as pH, is calculated based on the concentration of hydrogen ions (\(\text{H}^+\)) present. This measurement is fundamental to chemistry, as small changes in pH can profoundly affect chemical and biological reactions. All living systems, including the human body, rely on a narrow and stable pH range to ensure biological processes, like enzyme function, proceed correctly. If the pH deviates too far, it can lead to cellular damage and system failure, such as acidosis or alkalosis in the blood. To safeguard against these fluctuations, nature and science employ chemical compounds called buffers, with the phosphate buffer system being particularly important for maintaining internal stability.
What is a Buffer Solution
A buffer solution is a chemical system designed to resist changes in pH when a small amount of a strong acid or a strong base is added. Buffers are typically composed of a weak acid and its corresponding conjugate base existing in chemical equilibrium. The weak acid neutralizes any added base (\(\text{OH}^-\)), while the conjugate base neutralizes any added acid (\(\text{H}^+\)). This system absorbs excess ions, preventing a significant shift in the hydrogen ion concentration.
When a strong acid is introduced, the conjugate base reacts with the excess hydrogen ions, converting the strong acid into the buffer’s weak acid component. Conversely, when a strong base is introduced, the weak acid reacts with the added hydroxide ions, forming the conjugate base and water. By replacing the highly reactive strong acid or base with a much less reactive weak species, the buffer minimizes the change in pH. A buffer works best when the solution’s pH is centered around its \(\text{pKa}\) value, which is the point where the concentrations of the weak acid and conjugate base are equal.
The Chemistry of Phosphate Buffers
The phosphate buffer system utilizes components derived from phosphoric acid. The two main components at biological \(\text{pH}\) are dihydrogen phosphate (\(\text{H}_2\text{PO}_4^-\)) and hydrogen phosphate (\(\text{HPO}_4^{2-}\)), which form a conjugate acid-base pair. Dihydrogen phosphate acts as the weak acid, and hydrogen phosphate is its conjugate base. This pair is involved in a reversible chemical equilibrium that counteracts changes in \(\text{H}^+\) or \(\text{OH}^-\) concentration.
The \(\text{pKa}\) for this equilibrium is approximately 6.8 to 7.2, making the phosphate buffer highly effective in the \(\text{pH}\) range of 6.0 to 8.0. When a strong acid is added, the conjugate base (\(\text{HPO}_4^{2-}\)) accepts the excess hydrogen ions (\(\text{H}^+\)) to form more of the weak acid (\(\text{H}_2\text{PO}_4^-\)). This neutralizes the strong acid by converting it into a much weaker acid species.
If a strong base, such as hydroxide (\(\text{OH}^-\)), is introduced, the weak acid (\(\text{H}_2\text{PO}_4^-\)) donates a proton to the hydroxide, forming water (\(\text{H}_2\text{O}\)) and the conjugate base (\(\text{HPO}_4^{2-}\)). In both scenarios, the system shifts its equilibrium to consume the added strong ion, stabilizing the \(\text{pH}\). The proximity of the phosphate buffer’s \(\text{pKa}\) to the neutral \(\text{pH}\) of 7.0 makes it relevant for biological and laboratory environments.
Biological and Laboratory Applications
The phosphate buffer system plays a significant role in maintaining stable internal conditions within living organisms, particularly inside cells. It is a major buffering agent in the intracellular fluid, where phosphate ion concentration is high. This function is necessary because metabolic processes constantly produce small amounts of acid that must be neutralized. While the bicarbonate system dominates blood plasma, the phosphate buffer regulates \(\text{pH}\) inside the cell and in the kidney tubules.
In the kidneys, the phosphate buffer assists in the excretion of acid waste by binding to hydrogen ions, allowing them to be passed out in the urine. The lower \(\text{pH}\) of the fluid in the kidney tubules is closer to the buffer’s \(\text{pKa}\), maximizing its capacity to absorb \(\text{H}^+\) ions. This mechanism contributes to the body’s overall strategy to maintain systemic acid-base balance.
In laboratory and industrial settings, the phosphate buffer is widely used, often as Phosphate-Buffered Saline (PBS). PBS is a water-based salt solution formulated to match the \(\text{pH}\) (around 7.4) and osmotic pressure of human body fluids, making it isotonic and non-toxic to most cells. Researchers use PBS for various purposes, including washing cells and tissues, diluting substances, and maintaining a stable environment for cell culture experiments. Its stability and physiological compatibility also make it valuable in pharmaceutical applications.