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Acids and Bases

Organic Acids and Bases


One of the most important equations in chemistry (and for the MCAT) is the Henderson-Hasselbalch equation.

pH = pKa + log([conj. base]/[conj. acid]).

In Chem 101, often we learn to see the Henderson-Hasselbalch equation only 'straight-ahead'. We learn to look at it only as a function machine that tells us how the pH is a function of the relative predominance of acid base forms in a buffered solution.

However, in biochemistry especially, is it very useful to think of this equation the other way around, for predicting the ionization state of minor solution components. In the physiological context, substances with acid-base properties often exist within a pH environment that is determined by over-riding buffer systems. In that case, think of the Henderson-Hasselbalch equation the other way around. Think of the pH as already determined. The state of ionization of a minor component in the cytosol, such as a protein, would depend on the pH of its environment, determined by other physiological buffers present at much higher concentrations.

Suppose an ionizable group upon our protein has a pKa of 11, an amine group, for example. Which form is it in? Is the amine in the nonionized amine form (base) or the ammonium ion form (acid). If the solution is buffered to a stable pH of 8, the Henderson-Hasselbalch equation tells you that the ammonium ion form will predominate to a ratio of 1000 to 1. If the pKa of the molecule is less than the pH of its environment, the conjugate acid predominates.

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