Henderson-Hasselbach equation
Henderson-Hasselbalch Equation
The Henderson-Hasselbalch Equation is a mathematical formula that relates the pH of a buffer solution to the pKa (acid dissociation constant) of the acid and the ratio of the concentrations of the anion (A−) form of the acid to the undissociated acid (HA). It is a fundamental equation in the field of acid-base chemistry and is widely used in biology, chemistry, and medicine to calculate the pH of buffer solutions and to understand the buffering capacity of blood and other biological fluids.
Formula[edit | edit source]
The Henderson-Hasselbalch equation is given by:
- pH = pKa + log10([A−]/[HA])
where:
- pH is the hydrogen ion concentration of the solution,
- pKa is the acid dissociation constant, a measure of the strength of the acid,
- [A−] is the concentration of the anion form of the acid,
- [HA] is the concentration of the undissociated acid.
History[edit | edit source]
The equation was independently derived by two scientists, Lawrence Joseph Henderson in 1908 and Karl Albert Hasselbalch in 1916. Henderson first described the equation in the context of carbonic acid in blood. Hasselbalch later used the equation to study acid-base balance in blood, making it more widely known and applied in physiological contexts.
Applications[edit | edit source]
The Henderson-Hasselbalch equation is used in various scientific fields for different purposes:
- In biochemistry, it helps in understanding the buffering systems in biological fluids, such as the bicarbonate buffer system in blood.
- In pharmacy and medicine, it is used to calculate the pH of buffer solutions, which is crucial for drug formulation and stability.
- In environmental science, it aids in assessing the acidity and alkalinity of natural waters, which is important for aquatic life.
Limitations[edit | edit source]
While the Henderson-Hasselbalch equation is a useful tool, it has limitations. It assumes that the temperature is constant and that the activity coefficients of the ions are equal to one, which is not always the case in real solutions. Additionally, it is most accurate for buffer solutions where the concentrations of the acid and its conjugate base are similar and in the pH range of pKa ± 1.
See Also[edit | edit source]
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