Relationship of Charge to pH

pH is the negative log of the hydrogen ion concentration in a solution. Similarly pKA is the negative log of the acid dissociation constant for an amino acid side chain.

In solution, if pH < pKA, then the protonated form of an amino acid side chain predominates according to the Henderson-Hasselbalch equation . If pH > pKA, then the deprotonated form of the amino acid side chain predominates. The charge on an acidic side chain can therefore vary between -1 (when pH << pKA) and 0 ( when pH >> pKA). Consider the example of glutamate, with a side chain pKA = 4.4. When the pH = pKA, the concentration of the protonated and deprotonated forms is equal and the charge is -0.5. When pH = 6.4, the glutamate is about 99% deprotonated, leading to a charge of -0.99. When pH = 2.4, the glutamate is about 99% protonated, leading to a charge of -0.01.

The charge on a basic side chain can range between 0 (when pH >> pKA) and +1 (when pH << pKA). The side chain on lysine normally has a pKA value of about 9.0. When pH = pKA, the concentration of the protonated and deprotonated forms is equal and the charge is +0.5. When pH = 11.0, the lysine is about 99% deprotonated, leading to a charge of +0.01. When pH = 7.0, the lysine is about 99% protonated, leading to a charge of +0.99.

The graph below shows the relationship between charge and pH for the side chains of glutamate, an acidic amino acid, and lysine, a basic amino acid, assuming that these amino acids are incorporated in a protein and that they have pKA values of 4.4 for glutamate and 10.0 for lysine.

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