In chemistry, the Henderson Hasselbalch equation describes the derivation of pH as a measure of acidity (using pK_{a}, the negative log of the acid dissociation constant) in biological and chemical systems. The equation is also useful for estimating the pH of a buffer solution and finding the equilibrium pH in acidbase reactions (it is widely used to calculate the isoelectric point of proteins).
The Henderson Hasselbalch Equation is given by:
Here, [HA] is the molar concentration of the undissociated weak acid, [A⁻] is the molar concentration (molarity, M) of this acid’s conjugate base & pK_{a} is −log_{10}K_{a} where K_{a} is the acid dissociation constant, that is:

For the nonspecific Brønsted acidbase reaction:
In these equations, A⁻ denotes the ionic form of the relevant acid. Bracketed quantities such as [base] and [acid] denote the molar concentration of the quantity enclosed.
The HendersonHasselbalch equation is valid when it contains equilibrium concentrations of an acid and a conjugate base. In the case of solutions containing notsoweak acids (or notsoweak bases) equilibrium concentrations can be far from those predicted by the neutralization stoichiometry.
Here are several observations from Henderson Hasselbalch equation:
 If the pH = pK_{a}, the log of the ratio of dissociate acid and associated acid will be zero, so the concentrations of the two species will be the same. In other words, when the pH equals the pK_{a}, the acid will be half dissociated.
 As the pH increases or decreases by one unit relative to the pK_{a}, the ratio of the dissociate form to the associated form of the acid changes by factors of 10. That is, if the pH of a solution is 6 and the pK_{a} is 7, the ratio of [ A]/[ HA] will be 0.1, will if the pH were 5, the ratio would be 0.01 and if the pH were 7, the ratio would be 1.
 If the pH is below the pK_{a}, the ratio will be < 1, while if the pH is above the pK_{a}, the ratio will be >1.
 The HendersonHasselbalch equation can be also used in the case of polyprotic acids, as long as the consecutive pK_{a} values differ by at least 2 (better 3). Thus it can be safely used in the case of phosphoric buffers (pK_{a1}=2.148, pK_{a2}=7.199, pK_{a3}=12.35), but not in the case of citric acid (pK_{a1}=3.128, pK_{a2}=4.761, pK_{a3}=6.396).
Limitations
 The assumption that the concentration of the acid and its conjugate base at equilibrium will remain the same as the formal concentration. This neglects the dissociation of the acid and the binding of H+ to the base.
 The dissociation of water and relative water concentration itself is neglected as well.
These approximations will fail when dealing with relatively strong acids or bases (pKa more than a couple units away from 7), dilute or very concentrated solutions (less than 1 mM or greater than 1M), or heavily skewed acid/base ratios (more than 100 to 1). In high buffer dilutions, where the concentration of protons arising from water become equally or more prevalent than the buffer species themselves (at pH 7, this means buffer component concentrations of <10^{−5} M formally, but practically much higher), the pKa of the ‘buffer’ system will tend towards neutrality.
That’s all there’s to know about the equation. Want to know how Chemistry plays a role in our everyday life click here.
All the best!