ACID-BASE PHYSIOLOGY

Hydrogen Ions and pH

Small changes in the concentration of hydrogen ions can affect the ionic charge carried by enzymes, which can have a dramatic effect on their conformation and physiologic activity. Marked changes in hydrogen ion concentration can permanently denature enzymes and cause cell death. Thus, precise control of hydrogen ion concentration is essential.

By convention, pH (the negative log₁₀ of the hydrogen ion concentration), rather than the absolute concentration of hydrogen ions, is reported. The normal value for plasma pH is 7.4. An arterial pH below 7.35 constitutes an acidemia and a pH above 7.45 constitutes an alkalemia. Acidosis refers to a process whereby excess hydrogen ions are generated but, because of the body’s pH defense mechanisms, pH may not fall below 7.35. Alkalosis refers to a process whereby hydrogen ions are consumed but it does not necessarily result in a plasma arterial pH above 7.45.

Normal Regulation of pH

The body generates more than 200 liters of carbon dioxide and 80 mmol of nonvolatile acid each day as the result of oxidative metabolism. Despite this, the hydrogen ion concentration of extracellular fluid is kept relatively constant at approximately 0.000040 mmol/l. Thus, the pH of extracellular fluid is under constant threat from this large acid load. pH stability is achieved by three methods: (1) chemical buffers; (2) the lungs, which regulate the partial pressure of carbon dioxide; (3) the kidneys, which regulate the plasma bicarbonate concentration.

Buffers

A buffer is any compound that can reversibly bind with hydrogen ions and thereby attenuate the change in pH of a solution that arises from the addition of a hydrogen ion load.

All human buffers are weak acids. Weak acids are compounds that have a modest tendency to dissociate and release hydrogen ions (H⁺):

(31-1)

where HA is the weak acid and A⁻ is the conjugate base of that acid. The buffering solution contains both HA and A⁻ in equilibrium. If hydrogen ions are added to the solution, the equilibrium of this equation shifts to the left and some, but not all, of the hydrogen ions are removed from solution as HA. Thus, the fall in pH is attenuated. Conversely, if hydrogen ions are removed from solution, the equilibrium shifts to the right, and the rise in pH is minimized.

The ability of a buffer to attenuate a change in pH is determined by the tendency of the buffer to dissociate, which is characterized by the equilibrium constant (Ka) of the buffer:

(31-2)

Buffering ability is maximal when the concentration of the weak acid and its conjugate base are equal; that is, when [HA] = [A⁻], which occurs when the hydrogen ion concentration of the solution is equal to the dissociation constant (Ka) of the buffer; put another way, when the pH of the solution is equal to the pKa of the buffer (where pKa = the −log₁₀Ka). To function effectively, a buffer must have a pKa of within 1.0 unit of the pH of the solution being buffered.

Buffer systems may be open or closed. A closed buffer system is simply a fixed amount of buffer in solution. As hydrogen ions are added to the buffer system, [A⁻] falls, [HA] rises, and the ability of the buffer system to attenuate any further pH change diminishes. With an open buffer system, [HA] and [A⁻] can be independently controlled; for instance, HA can be eliminated from the solution and A⁻ regenerated. Thus, open buffer systems have a far greater capacity to deal with acid or alkali loads than do closed buffer systems.

Thus, the ability of a buffer to attenuate the change in pH of a solution in response to the addition or removal of hydrogen ions depends on: (1) the concentration of the buffer; (2) the starting pH of the solution; (3) the pKa of the buffer; and (4) whether the buffer system is open or closed.

Buffering Systems Within the Body.

There are four important buffers of extracellular fluid: (1) bicarbonate; (2) hemoglobin; (3) plasma proteins; (4) phosphate.

The bicarbonate buffer system is as follows:

(31-3)

Despite not having the same form as Equation 33-1, the bicarbonate buffer system reacts to changes in pH in a similar manner: the addition of H⁺ shifts the equation to the left, generating CO₂ and H₂O; the removal of H⁺ shifts the equation to the right, generating HCO₃⁻ (and H⁺). The pKa of the bicarbonate buffer system is 6.1, which at first glance seems too far removed from the pH of plasma to be very effective. However, because the concentration of bicarbonate in extracellular fluid is high (24 mmol/l) and because the bicarbonate buffer system is open, with the ability to independently control carbon dioxide and bicarbonate concentrations, it is quantitatively the most important buffer of the extracellular fluid.

The histidine amino acid found on hemoglobin and plasma proteins (pK_a 6.5) constitutes the majority of the nonbicarbonate buffering in the extracellular fluid. Although hemoglobin is an intracellular protein, the permeability of the red cell membrane is such that it can be considered an extracellular buffer.

Dihydrogen phosphate (pKa 6.8) is a relatively unimportant extracellular buffer because its concentration in plasma is low, approximately 1 mmol/l. However, dihydrogen phosphate plays an important role in buffering urine within the renal tubules, where it is concentrated. It is also an important buffer within cells.

Because all extracellular buffers are in solution with the same hydrogen ions, in effect, all buffers are in equilibrium with each other. This is known as the isohydric principle. Only the bicarbonate buffer system is open, so all of the extracellular buffers are regenerated by the elimination of carbon dioxide and the regeneration of bicarbonate. Furthermore, the pH of the entire system can be calculated from the Henderson-Hasselbalch equation as it applies to the bicarbonate buffer system:

(31-4)

where 6.1 is the pKa of the bicarbonate buffer; HCO₃⁻ is in mmol/l; and Pco₂ is in mmHg. The term 0.03 is the solubility coefficient of CO₂ in plasma and converts Pco₂ in mmHg into mmol/l of dissolved CO₂ or H₂CO₃.

In addition to extracellular buffers (and hemoglobin), some buffering occurs intracellularly. This is achieved by swapping hydrogen ions for potassium ions across cell membranes (to maintain electroneutrality). Intracellular buffering is quantitatively important in some forms of metabolic acidosis such as acute renal failure, and it contributes to the hyperkalemia seen in this condition.

Role of the Lungs: Control of P_aco₂

The carbon dioxide generated by oxidative metabolism is transported from the tissues to the lungs predominantly as bicarbonate (70%) or bound to hemoglobin (23%), with only a small proportion (7%) as dissolved gas. In the lungs, bicarbonate is reconverted to carbon dioxide and excreted in expired breath. Thus, the process of transporting carbon dioxide from the tissues to the lungs does not consume or generate bicarbonate buffer. However, the partial pressure of carbon dioxide in blood does influence pH.

The partial pressure of dissolved carbon dioxide in arterial blood (P_aco₂) is inversely proportional to alveolar ventilation and directly proportional to the rate of production of carbon dioxide (see Eq. 1-17). The carbon dioxide partial pressure in arterial blood is tightly controlled to a value close to 5.2 κPa (40 mmHg), corresponding to a concentration of dissolved gas of 1.2 mmol/l. A reduction in alveolar ventilation leads to an increase in P_aco₂, which drives the equilibrium of Equation 31-3 to the right, causing a fall in pH. This condition is termed respiratory acidosis. Conversely, an increase in alveolar ventilation drives the P_aco₂ down and the equilibrium of Equation 31-3 to the left, causing an increase in pH. This condition is termed respiratory alkalosis.

Role of the Kidneys: Control of Plasma HCO₃⁻ Concentration

The role of the kidneys in acid-base homeostasis is twofold: first, reabsorbing bicarbonate that is filtered at the glomerulus, and second, replacing bicarbonate that is consumed buffering metabolic acid.

On a normal diet, about 80 mmol/day of nonvolatile (non-carbon-dioxide or metabolic) acid is generated each day, mainly by the metabolism of proteins. This acid is buffered in the plasma by the bicarbonate buffer system and excreted as carbon dioxide gas, resulting in a net loss of bicarbonate. The cells of the distal convoluted tubule of the kidney regenerate this lost bicarbonate. In the process of regenerating bicarbonate, the distal tubular cells secrete an equimolar amount of hydrogen ions into the tubular fluid, resulting in acidification of the urine. With normal renal function, approximately 400 to 500 mmol of bicarbonate can be regenerated in this way (or, put another way, the kidneys can clear 400 to 500 mmol of H⁺ per day, providing that urinary buffering is ideal),^2,³ resulting in a urinary pH as low as 4.5. In this way, the kidneys regulate the plasma bicarbonate concentration (normal range 24 to 28 mmol/l) in the face of a changing metabolic acid load. The larger the metabolic acid load, the lower the urinary pH. This renal mechanism takes some hours to days to be fully effective.

From Equation 31-3 it is seen that an excess of hydrogen ions leads to a fall in plasma bicarbonate concentration, a condition known as metabolic acidosis. Conversely, the loss of hydrogen ions results in an increase in plasma bicarbonate concentration, a condition known as a metabolic alkalosis.