Sources of Acid

Acid production is derived from metabolism of sulfur containing (70 mEq/day), cationic   amino acids (140 mEq/day), and substances containing acid phosphates (H₂PO4^-, 30 mEq/day). Normally, this is balanced by H⁺ consumption in metabolism of anionic amino acids (110 mEq/day) and other organic anions (60 mEq/day), and by urinary excretion of ammonium (40 mEq/day) and acid phosphate (30 mEq/day) .

Buffering

Buffering occurs when pH changes are minimized. It involves chemical and physiological processes.
Chemical (in the ECF) and respiratory buffering are almost immediate. ICF buffering involves transfer of H⁺ in or out of the ICF.  Intracellular chemical buffering may be fast if it involves organic acid or slower (an intermediate step) for poorly cell permeable inorganic acids such as HCl or H₂SO₄. Long term regulation occurs through kidney urinary excretion of H⁺, with resulting reabsorption and generation of base (HCO₃^-).

Chemical buffering is due to conversion of a strong acid (such as HCl) into a weak acid (such as H₂CO₃ or H₂NaPO₄) through chemical reaction with a weak buffer base (such as NaHCO₃ or HNa₂PO₄). The higher the weak base concentration, the better the buffering. The weaker the resulting acid, the better the buffering.

The weakness or strength of an acid depends on its dissociation constant (K_d). A weak acid when dissolved yields few protons and its K_d is small. A strong acid when dissolved dissociates completely and yields many protons; its K_d is large. The pK is the negative log of the K_d, it is inversely related to the K_d:   pK= log (1/K_d). Weak acids have high pK. Strong acids have low pK.

Titration of a buffer

When an acid HA dissociates into H⁺ and buffer base (A^-), the mass action equilibrium is :

[H⁺] x [A^-] = [HA] x K_d  so that  [H⁺] = K_d x ([HA]/[A^-])

When [HA] = [A^-], then [H⁺] = K_d.

On a log scale, pH = pK + log ([HA]/[A^-]) and when [HA] = [A-], log [A-]/[HA]=log 1 = 0, so pH = pK.

When the pH of a buffered solution is more alkaline (higher) than the buffer pK by 1 unit, the log [A^-]/[HA] = 1and the ratio [A^-]/[HA] = 10. The proportion of the total buffer (HA + A^-) in the dissociated weak base form (A^-) will be 10/(10+1) = 0.91 (91%). When, by adding strong acid the solution, the pH becomes more acid than the buffer pK by one unit, the log [A^-]/[HA] = -1and the ratio [A-]/[HA] = 1/10. The proportion of the total buffer (HA + A^-) in the dissociated weak base form (A^-) will be 1/(10+1) = 0.091 (9.1%). So within +1 and -1 units of the buffer pK, 91-9.1 = 82% (most) of the buffer reacts (is titrated ) with H⁺ and is converted from A^- to HA. Beyond this pK region, there is little buffering of the pH on adding acid to the buffer containing solution. Thus, buffering is only effective within the pK region of the buffer (+/- 1 pH unit).

Isohydric principle.  When a solution (or compartment) contains more than one buffer, all buffer pairs (HA and A-) in the system are in equilibrium with the same proton concentration: Only those buffers with a pK within 1 pH unit of that in the solution participate effectively in the buffering of the solution pH.

Buffering power or capacity depends on the buffer concentrations and on the pK of the buffers present in that compartment.

Body buffers

The most abundant buffers in the body fluids are 1)  bicarbonate/CO₂ in the ECF (400 mEq), 2) bicarbonate in the ICF and bone (another 400 mEq), 3) histidine groups of intracellular proteins (about 400 mEq), and 4) small amounts of intra and extracellular phosphates (40 mEq). These buffers (total about 1250 mEq) maintain the ECF and ICF pH within 0.3-0.4 pH units of their normal values as long as their capacity to bind protons (about 20 mEq/Kg body weight) is not exceeded.