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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