Acid-base homeostasis

Neural structures in the regulation of acid-base balance

The central nervous system (CNS) is responsible for the total body content of CO ₂ , achieved as a result of the modulation of both voluntary and involuntary respiration.

• Voluntary respiration, as well as the hyperventilation seen in anxiety and certain primary CNS lesions, are determined by higher cortical centers.

• Involuntary respiration is controlled by areas in the brain stem and depends on input from both central and peripheral sensory receptors that respond to the concentrations of hydrogen ion, PCO ₂ , and partial pressure of oxygen (PO ₂ ) of the blood and cerebrospinal fluid (CSF).

The central component in involuntary ventilatory control is the respiratory control center in the medulla oblongata, which is composed of several nuclei. The medulla is responsible for establishing the basic respiratory pattern and integrates input from multiple sources, including higher brain centers, central and peripheral chemoreceptors, and baroreceptors.

Central chemoreceptors are distinct from the respiratory control center, but are located adjacent to it at the ventrolateral surface of the medulla.

• Two areas of chemosensation have been identified in this region, one caudal and one rostral, which sense changes in the pH and PCO ₂ of the brain stem interstitial fluid (ISF).

• Gap junctions in the pia may permit mixing of the brain ISF and CSF, allowing the chemosensitive areas to sense changes in the pH, PCO ₂ , and bicarbonate concentrations of the CSF.

• The central chemoreceptors are relatively insensitive to changes in PO ₂ except in cases of severe hypoxia.

Peripheral sensory input is contributed by pulmonary stretch receptors and carotid sinus and aortic arch chemo- and baroreceptors.

• The most important for maintenance of systemic acid-base balance is the carotid chemoreceptors.

• The carotid bodies, surrounded by a capillary plexus affording close proximity to systemic blood, respond to hypercapnia (high PCO ₂ ) and hydrogen ion concentrations.

• Their response to both PCO ₂ and to hydrogen ion concentration is virtually linear in the range of PCO ₂ from 25 to 65 mm Hg and a hydrogen ion concentration of 25 to 60 mEq/L.

• The carotid chemosensor is less sensitive to hypoxia with a PO ₂ of less than 60 mm Hg required for stimulation. Recall that at these PO ₂ levels, hemoglobin desaturation becomes significant.

• The carotid chemoreceptors’ sensitivity to combined hypoxia and hypercapnia exceeds the additive effect of the response to each stimulus individually: hypoxia renders the chemoreceptors more sensitive to hypercapnia, and vice versa.

An important concept in understanding the regulation of ventilation is that CO ₂ permeates into the CNS across the blood-brain barrier more quickly than HCO ₃ ^– . Therefore sudden increases in PCO ₂ peripherally will cause CO ₂ to enter and acidify the brain ISF and promote hyperventilation. A sudden decrease in PCO ₂ will allow CO ₂ to leave the CNS, giving a transient alkalinization of the brain ISF, depressing ventilation.

A review of chemical equilibria and buffers

When we first learn about chemical reactions, we hear about reactants and products, and we learn that some reactions can proceed in reverse as well as forward.

• Forward reaction between molecules of reactants occurs much more easily and frequently than the reverse reaction between molecules of products ( Fig. 22.1 ).

  

• The reverse reaction may still occur on a smaller scale and does so at the same time that the forward reaction proceeds. In this equation, the two-way arrow denotes reversibility.

  $\text{A}\text{+}\text{B}\leftrightarrow \text{C}\text{+}\text{D}$

Also recall that in first-order kinetics, the rate of the forward reaction is proportional to the concentration of the reactants in solution. Conversely, the rate of the reverse reaction is proportional to the concentration of the products.

• The more reactant molecules there are, the more forward reaction occurs per unit time.

• In a forward reaction that is energetically favorable, the reactants convert to products until the reactant concentrations have dropped considerably. In turn, the product concentrations have risen.

• The rate of the forward fast reaction decreases along with the dropping reactant concentration, and the rate of the reverse slow reaction climbs along with rising product concentration.

• When the forward and reverse rates meet and become equal, the concentrations of reactants and products can no longer change, and chemical equilibrium is reached.

Most metabolic reactions and drug metabolic processes occur with first-order kinetics, but recall that in some cases, the rate of product formation is independent of reactant concentration, and therefore linear (zero order, as in the case of alcohol metabolism when excess alcohol has saturated the enzyme alcohol dehydrogenase). In other cases, the reaction rate is related to the square of the reactant concentration (second-order kinetics).

If the energetics favor the forward reaction, the equilibrium is pushed “to the right,” toward the products.

• When the forward reaction is very energetically favorable, the product concentration must build up quite high before forward and reverse reactions have an equal rate and equilibrium is reached.

• The more favorable the forward reaction is, the larger the ratio of products to reactants at equilibrium will be.

• This equilibrium ratio of products to reactants is described by the equilibrium constant expression:

  $\text{K}\text{=}\left[\text{C}\right]\left[\text{D}\right]/\left[\text{A}\right]\left[\text{B}\right]$

K is the equilibrium constant, which reflects the ratio of forward to reverse rate constants of the reaction. The more favorable the forward reaction is, the higher K will be.

The opposite holds for a reaction in which the energetics are more favorable to the reverse reaction. Then the equilibrium is pushed “to the left,” toward the reactants.

• This implies that the K ratio of products to reactants at equilibrium must be lower.

• An irreversible reaction is one in which the reactants completely disappear as products appear; consequently, the K is infinite.

Adding to or subtracting from the concentrations of the reactants or products in equilibrium will cause a shift in the concentrations of reactants and products until they again reach equilibrium, where the ratio of products to reactants is again, K.

In Fig. 22.1 , if more A is added, for example, the forward reaction rate climbs, and more product is made.

• The forward rate then falls, and the reverse rate climbs, until equilibrium ratio K is again achieved.

• If D is subtracted (e.g., by reaction with another chemical species), the reverse reaction rate falls, and net product is made until the forward and reverse rates are again equal.

Equilibrium equations are also written for dissociation reactions, such as when an acid HA (considered the reactant) dissociates into products H ⁺ and an anionic conjugate base A ^– :

Acids are defined as strong acids or weak acids, depending on their tendency to dissociate into H ⁺ and A-. Stronger acids are more inclined toward dissociation and therefore have larger Ks, and weaker acids are less inclined toward dissociation and have smaller Ks.

Buffer solutions are solutions in which a weak acid (reactant) and conjugate base (product) are present together in similar concentrations at equilibrium.

• When a strong base is added, swallowing up protons, more net weak acid dissociates, producing more H ⁺ and A ^– until the equilibrium ratio of products to reactants is recovered, and the weak acid therefore acts as a reservoir for protons to restore the [H ⁺ ].

• When a strong acid is added, raising [H ⁺ ], the conjugate base associates with added free H ⁺ , producing more HA until recovery of the equilibrium ratio, and the conjugate base thus acts as a sponge for protons.

• In an unbuffered solution, adding acid will increase the free proton concentration dramatically, causing a steep drop in the pH. Adding base to an unbuffered solution will decrease the free proton concentration dramatically, causing a steep rise in pH.

It is important to understand that buffers lessen changes in pH but do not eradicate them.

The bicarbonate buffer in the blood

The most important extracellular buffer in body fluids is the bicarbonate/CO ₂ system.

The buffer is composed of a weak acid, carbonic acid (H ₂ CO ₃ ^– ), and a conjugate base, bicarbonate (HCO ₃ ^– ). Because of the rapid interconversion of CO ₂ and HCO ₃ ^– , we sometimes ignore the presence of H ₂ CO ₃ in the reaction sequence shown previously, and instead write:

The equilibrium equation is the following, with α equal to 0.03 mmol/L per mm Hg, the constant that accounts for dissolved CO ₂ and H ₂ CO ₃ in solution:

The bicarbonate buffer behaves in accordance with the description of buffers just shown.

• When protons are added, most of them will react with bicarbonate to form CO ₂ , as dictated by the buffer’s equilibrium constant.

• Proton losses will cause CO ₂ to combine with water, and then break down into HCO ₃ ^– and H ⁺ , thus restoring some of the lost H ⁺ and preventing a severe change in pH.

• Increases in bicarbonate consume some protons and raise the pH to a limited extent.

• Loss of bicarbonate will cause a shift to the right in the reaction. More CO ₂ will dissociate into H ⁺ and HCO ₃ ^– , and the increase in H ⁺ means a drop in pH. Because it is a buffer system, the equilibrium constant dictates that much of the acid (CO ₂ ) stays as CO ₂ , and therefore the pH drop is limited.

• Increases in CO ₂ will also shift the reaction to the right, elevating the free [H ⁺ ], but again, the equilibrium constant dictates that much of the added CO ₂ stay as CO ₂ .

• Decreases in CO ₂ cause protons and bicarbonate to react to replace the lost CO ₂ , raising the pH to a limited extent.

In considering the bicarbonate buffer relationships, physiologists use a more convenient form of the equilibrium equation, called the Henderson-Hasselbalch equation. It relates pH directly to the PCO ₂ and [HCO ₃ ^– ], mathematically summarizing the notion that increases in CO ₂ drop the pH (and vice versa) and decreases in [HCO ₃ ^– ] drop the pH (and vice versa):

Because the best buffers are those that have a pK close to the pH of body fluids (ratio of base-to-acid close to 1), one might expect the HCO ₃ ^– /CO ₂ buffer pair not to be a good buffer with its pK of 6.1, far from the normal pH of 7.4.

• The reason that HCO ₃ is considered a good buffer at physiologic pH, is because the system works in an “open” fashion in which CO ₂ is allowed to escape from the system via ventilation (as opposed to a “closed” container).

• The lungs prevent buildup of acid and CO ₂ when an excess of H ⁺ is added to the body.

Other buffers in the body

In addition to the bicarbonate buffer system, there are other buffers in the ECF and the ICF. The same body fluid H ⁺ concentration is simultaneously in equilibrium with multiple buffer pairs, the importance of each depending on their pK and the amount of buffer present.

This isohydric principle allows for the acid-base status to be examined by close evaluation of a single buffer pair, such as HCO ₃ ^– /PCO ₂ .

• Protein and phosphates both act as weak acid buffers inside and outside of cells.

• Buffering inside cells by protein and phosphate takes place when an increase in the extracellular [H ⁺ ] leads to the diffusion of protons into cells. (In red blood cells, hemoglobin is a protein that buffers [H ⁺ ] changes intracellularly.)

Intracellular buffering may account for half or more than half of the buffering of acid loads

• When increased plasma [H ⁺ ] shifts from the ECF to the intra-cellular fluid (ICF), important changes occur in the levels of other cations in the blood.

• When the protons diffuse into the cell, there is an increase in the positive charge inside the cell, slightly depolarizing the membrane, and thus reducing the electrical gradient that opposes K ⁺ efflux from the cell.

• The net result is that H ⁺ enters the cell and K ⁺ leaves, raising the plasma [K ⁺ ].

• This is why acidemia (low serum pH) is often associated with hyperkalemia (high serum [K ⁺ ]), as discussed in Chapter 21 .

Acid loads

Metabolically generated acid can be divided into two main groups:

Gaseous or volatile acid (CO ₂ ) is produced by oxidative metabolism of carbohydrates and fats, on the order of 15,000 to 20,000 mmol/day.

• The production of CO ₂ raises PCO ₂ , raises [H ⁺ ], and lowers blood pH, in accordance with the equilibrium equation described previously.

• The kidneys respond to the acidity, and the lungs respond to the increased CO ₂ (and to the acidity) to counter the drop in pH.

• CO ₂ accounts for the vast majority of acid produced by the body (see Fast Fact Box 22.2 ).

The metabolism of amino acids, nucleic acids, and other compounds releases acidic and alkaline byproducts that are not gaseous (nonvolatile), sometimes called fixed acids and bases.

• The breakdown of sulfur-containing amino acids, such as cysteine and methionine, yields acid sulfates, whereas lysine, arginine, and histidine are often hydrochlorides.

• The digestion of organic phosphorus-containing compounds gives rise to acid phosphates.

• Alkaline byproducts come from the metabolism of anionic amino acids, such as aspartate and glutamate, with accompanying strong cations, such as Na ⁺ , K ⁺ , and Ca ^{2 +} .

• The normal, high-protein American diet yields a greater quantity of acid than base. On average, a net of 1 mEq/kg per day of nonvolatile organic and inorganic acids is produced.

Acid loads are buffered, but they still acidify the blood to some extent. The addition of H ⁺ to the blood raises [H ⁺ ], lowers pH, and raises PCO ₂ in accordance with the equilibrium equation. The kidneys respond to the acidity, and the lungs respond to the increased CO ₂ (and to the acidity) to counter the drop in pH.

The reabsorption of filtered HCO ₃ ^–

Before we consider the kidney’s response to acid loads, we should consider the consequences of the fact that bicarbonate is freely filtered across the glomerulus.

The daily filtered HCO ₃ ^– totaling about 4320 mEq of HCO ₃ ^– (the product of a glomerular filtration rate of 180 L/day and an [HCO ₃ ^– ] of 24 mEq/L) is presented to the tubules.

• Clearly, the first critical step confronting the tubule is to reabsorb this enormous filtered quantity, because loss of even a small fraction of that amount would result in severe acidosis.

• Therefore the kidney not only responds to acid loads by creating new bicarbonate and excreting acid, but must also first reabsorb all of the bicarbonate in the tubular filtrate to maintain the blood pH at normal levels.

How does the reabsorption of bicarbonate take place?

• Renal tubule cells do not have any apical transporters to reabsorb bicarbonate directly.

• Instead, the cells along the nephron use a more complex series of reactions to reabsorb the bicarbonate ( Fig. 22.2 ).

  

• Although the actual reabsorption reactions are more involved than direct transport of bicarbonate, the end result is that bicarbonate is translocated from the tubule lumen into the bloodstream at the expense of adenosine triphosphate (ATP).