A buffer is a solution with a weak acid and conjufate base in equilibrium. It resists changes in pH due to addition of other acids or bases.
Buffering allows for large changes in acid/base load with near-constant [H+].
Buffer systems can be
- Closed (total quantity [A-] + [AH] is constant), or
- Open (one/both of [A-] or [AH] can be excreted or generated \(\to\) shifts equilibrium by mass action \(\to\) physiological control of pH)
Buffering capacity is maximized when buffer concentration is high, the buffer is open, and the pH is near the buffer's pKa:
$$pH = pKa + \log \frac{[A-]}{[AH]}$$
All buffer systems share the same pool of hydrogen, so they are in equilibrium with one another (the "isohydric principle").
Bicarbonate: the major buffer of the ECF \(CO_2 + H_2O \rightleftharpoons H_2CO_3 \rightleftharpoons HCO_3^- + H^+\)
- pKa 6.1
- High concentration
- Both ends open (CO2 \(\propto \ V_A\) and kidneys can generate/secrete bicarbonate \(\to\) most important buffer system)
- Allows for rapid control of pH
Haemoglobin: the secondary buffer of the ECF
- OxyHb pKa 6.6, DeoxyHb pKa 8.2 (\(\pm\)0.8 from physiological 7.4)
- Very high concentration
- Therefore high buffering capacity
- Hydrogen ions bind to imidazole residues
- Increased propensity for DeoxyHb to bind hydrogen gives rise to the Haldane effect (\(\uparrow\)H+ binding \(\to\) equilibrium shifted to right, \(\downarrow\)PaCO2 \(\uparrow\)[HCO3-] \(\to\) Hamburger effect, HCO3- antiported out of cell, chloride in.
Intracellular proteins
- pKa 6.8
- Very high concentration
- H+ binds to imidazole residues
- Most important buffer of ICF
Phosphate: buffer of urine and ICF \(H_2PO_4^- \rightleftharpoons HPO_4^- + H^+\)
- pKa 6.1, low extracellular quantities
- Mostly important as intracellular buffer and buffer of urine
- Technically open, but slow. In chronic acidosis, active release of HPO4 from bone \(\to\) H2PO4 in renal tubule poorly absorbed \(\to\) lost in urine.