Biochemistry/pKa values

Biochemistry

<< Metabolism and energy

It is difficult to discuss the subject of biochemistry without a firm foundation in general chemistry and organic chemistry, but if one doesn't remember the concept of the Acid Dissociation Constant (pKa) from Organic Chemistry, one can read up on the topic below.

Buffers are essential to biochemical reactions, as they provide a (more or less) stable pH value for reactions to take place under constantly changing circumstances. The pH value in living cells tends to fall between 7.2 and 7.4, and this pH level is generally maintained by weak acids. (The pH values in lysosomes and peroxisomes differ from this value, as do the pH measurements of the stomach and other organs found in various types of plants and animals.)

An acid is here defined simply as any molecule that can release a proton (H+) into a solution. Stronger acids are more likely to release a proton, due to their atomic and molecular properties. The tendency of an acid to release a proton is called the dissociation constant (Ka) of that substance, with
 * $${K}_{a} = {[{H}^{+}][{A}^{-}] \over [HA]}$$

for HA <-> H+ + A-.

A larger Ka value means a greater tendency to dissociate a proton, and thus it means the substance is a stronger acid. The pH at which 50% of the protons are dissociated can therefore be calculated as:
 * pKa = -log ( Ka )

This equation is known as the Henderson-Hasselbalch equation.


 * $${pH} = {pK}_{a} + \log{[\text{base}] \over [\text{acid}]}$$

The Henderson-Hasselbalch equation is derived from the adjacent Ka expression. By taking the logarithm of base ten to both sides, the next part of the equation is obtained. Using the logarithmic property of multiplication, the [H+] breaks from the expression. Since log Ka is equal to -pKa and log [H+] is equal to -pH, they are then substituted. To obtain what is known as the Henderson-Hasselbalch equation, -pKa and -pH are subtracted from their respective sides to yield a positive equation.

The Henderson-Hasselbalch equation interestingly enough predicts the behavior of buffer solutions. A solution of 1 M ethanoic (acetic) acid [HA] and 1 M sodium ethanoate [A-] will have a pH equal to the pKa of ethanoic acid: 4.76.

If we added acid to pure water up to a concentration of 0.1 M, the pH would become 1. If we add the same concentration of acid to the buffer solution, it will react with the ethanoate to form ethanoic acid. The ethanoate concentration would drop to 0.9 M and the ethanoic acid concentration should rise to 1.1 M. The pH becomes 4.67 - very different from the pH=1 without a buffer.

Similarly, adding 0.1 M alkali changes the pH of the buffer to 4.85, instead of pH=13 without buffer. Due to the amphipathic nature of amino acids - which are the monomer building blocks of all proteins, physiological conditions are always considered to be buffered, which plays a major role in the conformations and reactivities of substrates in the cell's liquid interior, its cytosol.

A very small (which would include a large negative value) pKa indicates a very strong acid. A pKa value between 4 and 5 is the most common range for organic acid compounds.

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