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).

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

All the H⁺ that is buffered must eventually be excreted from the body via the kidneys, regenerating the bicarbonate used up in the buffering process and maintaining the plasma bicarbonate concentration within normal limits (Fig 20.2). Secretion of H⁺ by the tubular cells serves initially to reclaim bicarbonate from the glomerular filtrate so that it is not lost from the body. When all the bicarbonate has been recovered, any deficit due to the buffering process is regenerated. The mechanisms for bicarbonate recovery and for bicarbonate regeneration are very similar and are sometimes confused (Fig 20.2).

The excreted H⁺ must be buffered in urine or the [H⁺] would rise to very high levels. Phosphate acts as one such buffer, while ammonia is another (Fig 20.3).

Assessing status

Normal acid–base balance involves the bicarbonate buffer system. In chemical terms, the bicarbonate buffer system can be considered in the same way as any other chemical dissociation.

By the law of mass action:

where K is the first dissociation constant of carbonic acid.

But the carbonic acid (H₂CO₃) component is proportional to the dissolved carbon dioxide, which is in turn proportional to the partial pressure of the CO₂.

[H₂CO₃] can, therefore, be replaced in the mass action equation by PCO₂. At this point, an understanding of the role of the bicarbonate buffer system in assessing clinical acid–base disorders can be achieved simply by reference to the relationship:

which shows that the H⁺ concentration in blood varies as the bicarbonate concentration and PCO₂ change. If everything else remains constant:

Blood [H⁺] is controlled by our normal pattern of respiration and the functioning of our kidneys.

The [H⁺] is 40 nmol/L and [HCO₃⁻] is 25 mmol/L, i.e. 25 000 000 nmol/L. Thus, changes in their respective concentrations are not directly linearly comparable.

The acid–base status of the patient and the magnitude of the disturbance can be obtained by measuring the components of the bicarbonate buffer system.

Acid–base disorders

‘Metabolic’ acid–base disorders are those that directly cause a change in the bicarbonate concentration. Examples include diabetes mellitus, where altered intermediary metabolism in the absence of insulin causes a build up of H⁺ from the ionization of acetoacetic and β-hydroxybutyric acids, or loss of bicarbonate from the extracellular fluid, e.g. from a duodenal fistula.

‘Respiratory’ acid–base disorders affect directly the PCO₂. Impaired respiratory function causes a build-up of CO₂ in blood, whereas, less commonly, hyperventilation can cause a decreased PCO₂.