Acid–base equilibrium

The balance between acids (elements or compounds that increase H+ concentration) and bases (elements or compounds that decrease H+ concentration). Neutrality (equal balance of acid and base) is at a pH of 7.0 (H+ concentration = 1 10 7 mol l 1 ). However, homeostatic mechanisms in living organisms tend to maintain an extracellular fluid pH between 7.35 and 7.45. Survival of the organism is not possible outside of the range of a pH between 7.0 and 7.7. Acidosis is defined as a blood pH < 7.35 and occurs with prolonged starvat tion, severe diarrhoea, asphyxia, ketosis and lactic acidosis. Alkalosis is defined as a blood pH > 7.45 and is associated with hyperventilation, vomiting of gastric acid and diuresis. Three systems within the body are primarily responsible for maintenance and regulation of acid–base equilibrium. These are the physiological buffers, the respiratory system and the renal system. These systems are interrelated and provide relatively rapid responses to shifts in acid–base equilibrium. The gastrointestinal tract also plays important roles in acid–base equilibrium but the responses are of greater consequence to long-term regulation and involve shifts in absorption and excretion of mineral ions.

     Major physiological buffers include bicarbonate, phosphate and proteins. Bicarbonate

ions (HCO3 – ) and hydrogen ions (H+) are in equilibrium with carbonic acid (H2 CO3 ), a

weak acid. Carbonic acid is produced by enzymatic action of carbonic anhydrase from CO2 and H2 O. The formation and end-products of bicarbonate can be easily eliminated via respiratory or renal systems without an effect on pH. Since mechanisms exist to maintain a constant extracellular concentration of bicarbonate ions (which are an excellent buffer for physiological fluids), the bicarbonate buffer does not provide a means for net elimination of acidic or basic loads imposed on the body. In terms of acid–base equilibrium, the bicarbonate buffer is considered a futile cycle since net elimination of bicarbonate as CO2 via the lungs is eventually compensated for by renal synthesis of bicarbonate by the kidneys with no net change in H+. Phosphate ions buffer H+ in physiological fluids and contribute to the net equilibrium of acids and bases in the body. Within physiological pH ranges the concentration of dibasic (HPO4 – ) phosphate ions is

approximately four times the concentration of moanobasic (H2 PO4 – ), but the kidneys can concentrate H+ in urine to a pH as low as 4.5. As urine pH decreases, the dibasic phosphate ions provide a buffer by accepting H+ to form monobasic phosphate, thus providing net elimination of H+ from the body.

     Another major route for a net elimination of H+ from the body involves renal production and secretion of ammonium ions from glutamine catabolism. Under acid loads a trans porter in renal mitochondria is inhibited, resulting in additional degradation of glutamine and excretion of H+ as ammonium (NH4 +).

     The strong ion difference (SID), which is the sum of all strong cations (mol l 1 ) minus

the sum of all strong anions (mol l 1 ), also impacts on the regulation of acid–base equilibrium. The SID affects the partial pressure of blood CO2 and renal electrolyte excretion. Shifts in SID impact renal compensation by changes in the relative amounts of ammonium and phosphate ion excretion.