Review of acid-base concepts:

a. Water dissociates: 2 H₂O <==> H+-H₂O + OH-

b. Keq = [H⁺][OH^-]/[H₂O] Assume activity of water is 1.00, yields --->

c. Kw = [H⁺][OH^-] = 1.0 x 10^-14 at 25 C; Note the reciprocal relationship.

Definitions of acids and bases:

a. Bronsted-Lowry theory: Good for aqueous solutions and practical: acid = proton donor; base = proton acceptor

b. Lewis theory: More useful for organic reactions and more comprehensive Base = electron donor Acid = electron acceptor

Definition of pH:

pH = -log[H⁺] = log 1/[H⁺] Remember: 0 - log[H] = log 1/[H]

Definition of dissociation:

HA <==> H⁺ + A^- Kd (or K_a) = [H+][A^-]/[HA] K_a expresses the tendency of an acid to release protons (i.e. the larger the Ka the stronger the acid (in water)).

pK_a = -logK_a = log (1/K_a); pK_a = pH of half-dissociation

Henderson-Hasselbalch Equation:

This equation relates the pH, pK_a, and degree of dissociation

Derivation:

K_a = [H][A]/[HA] Remember that pK_a = log (1/K_a) and pH = log (1/[H]) **divide by [H] and K

1/[H] = (1/Ka)[A]/[HA] Take log of both sides

log (1/[H]) = log (1/K_a) + log([A]/[HA])

pH = pKa_a + log ([A]/[HA]) = HH Equation (or use log ([UP]/[P]); Brabson)

UP = unprotonated; P = protonated (works for both acids and bases!)

This is only one mathematical expression relating pH and pK, but it is frequently used by biochemists because it has an isolated term relating the ratio of charged and uncharged species.

Why is pH important in biochemistry?

1. Enzyme active site residues must be in the protonated or unprotonated form for full activity.

2. Need to control pH of medium to regulate this state and to maintain proteins in the active configuration (external side chain residues at proper state of ionization).

3. In some cases the proton is actually a reactant or product in the biochemical reaction.