The term pH comes from the French puissance hydrogen, meaning the strength or power of hydrogen. The hydrogen ion concentration of blood expressed in gram molecular weights of hydrogen per liter (moles/L) is so much less than 1 (e.g., 0.0000001) that it is easier to communicate this information in terms of logarithms; thus, 0.0000001 becomes 1 x 10–7. The symbol pH represents an even greater simplification, because pH is defined as the negative logarithm of the hydrogen ion concentration (in the preceding example, 1 x 10–7 becomes 10–7, which then becomes 7.0). In this way, a relatively simple scale is substituted for very cumbersome tiny numbers. In the pH scale, therefore, a change in 1.0 pH unit means a tenfold change in hydrogen ion concentration (a change of pH from 7.0 to 6.0 represents a change from 10–7 to 10–6 moles/L).

The normal pH of arterial blood is 7.4, with a normal range between 7.35 and 7.45. Blood pH must be maintained within relatively narrow limits, because a pH outside the range 6.8-7.8 is incompatible with life. Therefore, the hydrogen ion content is regulated by a series of buffers. A buffer is a substance that can bind hydrogen ions to a certain extent without inducing a marked change in pH. Among substances that act as buffers are hemoglobin, plasma protein, phosphates, and the bicarbonate-carbonic acid system. Bicarbonate is by far the body’s most important buffer substance; it is present in large quantities and can be controlled by the lungs and kidneys.

A brief review of the bicarbonate-carbonic acid system recalls that carbon dioxide (CO2) in aqueous solutions exists in potential equilibrium with carbonic acid (CO2 + H2O = H2CO3). The enzyme carbonic anhydrase catalyzes this reaction toward attainment of equilibrium; otherwise the rate of reaction would be minimal. CO2 is produced by cellular metabolism and is released into the bloodstream. There, most of it diffuses into the red blood cells (RBCs), where carbonic anhydrase catalyzes its hydration to H2CO3. Carbonic acid readily dissociates into hydrogen ions (H+) and bicarbonate ions (HCO–3). Eventually only a small amount of dissolved CO2 and a much smaller amount of undissociated H2CO3 remain in the plasma. Therefore, the great bulk of the original CO2 (amounting to 75%) is carried in the blood as bicarbonate, with only about 5% still in solution (as dissolved CO2 or undissociated H2CO3) and about 20% coupled with hemoglobin as a carbamino compound or, to a much lesser extent, coupled with other buffers, such as plasma proteins.

This situation can best be visualized by means of the Henderson-Hasselbalch equation. This equation states that    pH = pK + log(base/acid)  where pK is the dissociation constant (ability to release hydrogen ions) of the particular acid chosen, such as H2CO3. The derivation of this equation will be disregarded to concentrate on the clinically useful parts, the relationship of pH, base, and acid. If the bicarbonate-acid system is to be interpreted by means of the Henderson-Hasselbalch equation, then base/acid = HCO3- / H2CO3 since bicarbonate is the base and carbonic acid is the acid. Actually, most of the carbonic acid represented in the equation is dissolved CO2 which is present in plasma in an amount greater than 100 times the quantity of undissociated H2CO3. Therefore the formula should really be … but it is customary to let H2CO3 stand for the entire denominator. The next step is to note from the Henderson-Hasselbalch equation that pH is proportional to base/acid. This means that in the bicarbonate-carbonic acid system, pH is proportional to …  The kidney is the main regulator of HCO–3 production (the numerator of the equation) and the lungs primarily control CO2 excretion (the denominator).

The kidney has several means of excreting hydrogen ions. One is the conversion of monohydrogen phosphate to dihydrogen phosphate (HPO42– to H2PO4–). Another is the formation of ammonia (NH3) in renal tubule cells by deamination of certain amino acids, such as glutamine. Ammonia then diffuses into the urine, where it combines with H+ to form ammonium ion (NH4+). A third mechanism is the one of most concern now, the production of HCO–3 in the renal tubule cells. These cells contain carbonic anhydrase, which catalyzes the formation of H2CO3 from CO2. The H2CO3 dissociates, leaving HCO–3 and H+. The H+ is excreted in the urine by the phosphate or ammonium pathways or combined with some other anion. The HCO–3 goes into the bloodstream where it forms part of the pH buffer system. Not only does HCO–3 assist in buffering H+ within body fluids, but HCO–3 filtered at the glomerulus into the urine can itself combine with urinary H+ (produced by the renal tubule cells from the H2CO3 cycle and excreted into the urine).

The lungs, on the other hand, convert H2CO3, into CO2 and water with the aid of carbonic anhydrase and blow off the CO2. In this process, a mechanism for excreting H+ exists, because the HCO–3 in the plasma can combine with free H+ to form H2CO3, which can then be eliminated from the lungs in the form of CO2, as was just described. This process can probably handle excretion of most normal and mildly abnormal H+ quantities. However, when large excesses of free H+ are present in body fluids, the kidney plays a major role, because not only is H+ excreted directly in the urine, but HCO–3 is formed, which helps buffer additional amounts of H+ in the plasma.

Going back to the Henderson-Hasselbalch equation, which is now modified to indicate that pH is proportional to HCO–3/H2CO3, it is easy to see that variations in either the numerator or denominator will change pH. If HCO–3 is increased without a corresponding increase in H2CO3, the ratio will be increased and the pH will rise. Conversely, if something happens to increase H2CO3 or dissolved CO2, the denominator will be increased, the ratio will be decreased, the pH will fall, and so on. Clinically, an increase in normal plasma pH is called “alkalosis” and a decrease is called “acidosis.” The normal ratio of HCO–3 to H2CO3 is 20:1.