Acid–base: concepts and vocabulary

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Acid–base

concepts and vocabulary

H⁺ concentration

Blood hydrogen ion concentration [H⁺] is maintained within tight limits in health. Normal levels lie between 35 and 45 nmol/L. Values greater than 120 nmol/L or less than 20 nmol/L require urgent treatment; if sustained they are usually incompatible with life. The [H⁺] in blood may also be expressed in pH units. pH is defined as the negative log of the hydrogen ion concentration. (Fig 20.1).

Fig 20.1 The negative logarithmic relationship between [H⁺] and pH.

H⁺ production

Hydrogen ions are produced in the body as a result of metabolism, particularly from the oxidation of the sulphur-containing amino acids of protein ingested as food. The total amount of H⁺ produced each day in this way is of the order of 60 mmol. If all of this were to be diluted in the extracellular fluid (≈ 14 L), [H⁺] would be 4 mmol/L, or 100 000 times more acid than normal! This just does not happen, as all the H⁺ produced are efficiently excreted in urine. Everyone who eats a diet rich in animal protein passes a urine that is profoundly acidic.

Metabolism also produces CO₂. In solution this gas forms a weak acid. Large amounts of CO₂ are produced by cellular activity each day with the potential to upset acid–base balance, but under normal circumstances all of this CO₂ is excreted via the lungs, having been transported in the blood. Only when respiratory function is impaired do problems occur.

Buffering

A buffer is a solution of a weak acid and its salt (or a weak base and its salt) that is able to bind H⁺ and therefore resist changes in pH. Buffering does not remove H⁺ from the body. Rather, buffers temporarily mop up any excess H⁺ that are produced, in the same way that a sponge soaks up water. Buffering is only a short-term solution to the problem of excess H⁺. Ultimately, the body must get rid of the H⁺ by renal excretion.

The body contains a number of buffers to even out sudden changes in H⁺ production. Proteins can act as buffers, and the haemoglobin in the erythrocytes has a high capacity for binding H⁺. In the ECF, bicarbonate buffer is the most important. In this buffer system, bicarbonate (HCO₃⁻) combines with H⁺ to form carbonic acid (H₂CO₃). This buffer system is unique in that the (H₂CO₃) can dissociate to water and carbon dioxide.

Whereas simple buffers rapidly become ineffective as the association of the H⁺ and the anion of the weak acid reaches equilibrium, the bicarbonate system keeps working because the carbonic acid is removed as carbon dioxide. The limit to the effectiveness of the bicarbonate system is the initial concentration of bicarbonate. Only when all the bicarbonate is used up does the system have no further buffering capacity. The acid–base status of patients is assessed by consideration of the bicarbonate system of plasma.

The association of H⁺ with bicarbonate occurs rapidly, but the breakdown of carbonic acid to carbon dioxide and water happens relatively slowly. The reaction is accelerated by an enzyme, carbonic anhydrase, which is present particularly where this reaction is most needed, in the erythrocytes and in the kidneys. Buffering by the bicarbonate system effectively removes H⁺ from the ECF at the expense of bicarbonate. The carbon dioxide that is formed can be blown off in the lungs, and the water mixes with the large body water pool. The extracellular fluid contains a large amount of bicarbonate, around 24 mmol/L. If H⁺ begins to build up for any reason, the bicarbonate concentration falls as the buffering system comes into play.

H⁺ excretion in the kidney

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