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Chapter 14. Fluid, Electrolyte, and Acid-Base Balance

14.3 Acid-Base Balance

Learning Objectives

By the end of this section, you will be able to:

  • list the three main buffer systems in the body;
  • identify the most rapid buffer system in the body;
  • state the most important buffer in the ECF;
  • state the most important buffer in the ICF;
  • define compensation;
  • explain the how the respiratory system can regulate plasma pH; and
  • explain the how the kidneys can regulate plasma pH.

Maintenance of plasma pH within the homeostatic range, 7.35 to 7.45, is essential for healthy physiological function. Specifically, pH affects the functioning of all proteins; for example:

  1. All enzymes function optimally in a specific pH range; therefore, all metabolic pathways can be affected by pH changes.
  2. Hemoglobin function is affected by pH: Hb-O2 affinity decreases as pH decreases.
  3. Membrane proteins are affected by pH because a change in pH could affect the conformation of the protein.

Plasma pH is maintained by the following three mechanisms:

  1. buffer systems: these are the first to respond to an acid-base disturbance and act within seconds;
  2. respiratory mechanisms: these are the next to respond to an acid-base disturbance and can have an effect within 1 to 3 minutes;
  3. renal mechanisms: these are the last to respond to an acid-base disturbance, having an effect between 1 to 3 days, but are the most potent mechanism for correcting an acid-base disturbance.

Acidity of a solution is measured using the pH scale, as shown in Figure 14.3.1. Blood, at 7.35 to 7.45, is slightly alkaline. A variety of buffering systems permits blood and other bodily fluids to maintain a narrow pH range, even in the face of perturbations. A buffer is a chemical system that prevents a drastic change in fluid pH by dampening the change in H+ concentration in the case of excess acid or base. Most commonly, the buffering substance is either a weak acid, which takes up hydroxyl ions (OH–), or a weak base, which takes up H+.

This table gives examples of solutions from PH of zero to 14. Examples of solutions with a PH of zero include battery acid and strong hydrofluoric acid. An example of a solution with a pH of one is the hydrochloric acid secreted by the stomach lining. Examples of solutions with a PH of two include lemon juice and vinegar. Examples of solutions with a PH of three include grapefruit juice, orange juice and soda. Examples of solutions with a PH of four include tomato juice and acid rain. Examples of solutions with a PH of five include soft drinking water and black coffee. Examples of solutions with a PH of six include urine and saliva. An example of a solution with a PH of seven is pure water. An example of a solution with a PH of eight is sea water. An example of a solution with a PH of nine is baking soda. Examples of solutions with a PH of ten include saline lake water and milk of magnesia. An example of a solution with a PH of eleven is an ammonia solution. An example of a solution with a PH of twelve is soapy water. Examples of solutions with a PH of thirteen include bleach and oven cleaner. An example of a solution with a PH of fourteen is liquid drain cleaner.
Figure 14.3.1 – The pH Scale: Solutions with a pH less than 7.0 are acidic, and those with a pH greater than 7.0 are alkaline.

Buffer Systems in the Body

The buffer systems in the human body are extremely efficient, and different systems work at different rates. It takes only seconds for the chemical buffers in the blood to make adjustments to pH. The respiratory tract can adjust the blood pH in minutes by changing ventilation rate. The kidneys can adjust blood pH via changing the secretion of H+ and the reabsorption of bicarbonate, but this process takes hours to days to have an effect.

The buffer systems functioning in blood plasma include plasma proteins, phosphate, and the carbonic acid-bicarbonate buffer system.

Protein Buffers in Plasma and Cells

Nearly all proteins can function as buffers. Proteins are made up of amino acids, which contain positively charged amino groups and negatively charged carboxyl groups. The charged regions of these molecules can bind H+ and OH– ions, and thus function as buffers.

Protein buffers are the main buffer system in the ICF. In the ECF, the plasma proteins contribute to the buffering capability of the blood.

Hemoglobin as a Buffer

Hemoglobin is the principal protein inside of red blood cells and accounts for one-third of the mass of the cell. During the conversion of CO2 into bicarbonate, H+ liberated in the reaction are buffered by hemoglobin, which is reduced by the dissociation of oxygen. This buffering helps maintain normal erythrocyte pH.

Phosphate Buffer

Phosphates are found in the blood in two forms: sodium dihydrogen phosphate (Na2H2PO4−), which is a weak acid, and sodium monohydrogen phosphate (Na2HPO42-), which is a weak base. When Na2HPO42- comes into contact with a strong acid, such as HCl, the base picks up a second hydrogen ion to form the weak acid Na2H2PO4− and sodium chloride, NaCl:

HCl + Na2HPO4→NaH2PO4 + NaCl
(strong acid) + (weak base) → (weak acid) + (salt)

When Na2HPO42− (the weak acid) comes into contact with a strong base, such as sodium hydroxide (NaOH), the weak acid reverts back to the weak base and produces water.

NaOH + NaH2PO4→Na2HPO4 + H2O
(strong base) + (weak acid) → (weak base) + (water)

Phosphates are important buffers in urine and the ICF. Phosphates are low in concentration in the ECF, so are not important ECF buffers.

Carbonic Acid-Bicarbonate Buffer System

The carbonic acid-bicarbonate buffer system, discussed in the previous section of this chapter and in other chapters, is the most important ECF buffer. It is a mixture of carbonic acid and sodium bicarbonate, a weak base, and it works in similar fashion to the phosphate buffer system.

When sodium bicarbonate (NaHCO3), comes into contact with a strong acid, such as HCl, carbonic acid (H2CO3), which is a weak acid, and NaCl are formed:

NaHCO3 + HCl → H2CO3+NaCl
(sodium bicarbonate) + (strong acid) → (weak acid) + (salt)
When carbonic acid comes into contact with a strong base, such as NaOH, bicarbonate and water are formed:
H2CO3 + NaOH→HCO3- + H2O
(weak acid) + (strong base)→(bicarbonate) + (water)
As with the phosphate buffer, a weak acid or weak base captures the free H+ or OH-, and a significant change in pH is prevented.
Bicarbonate ions and carbonic acid are present in the blood in a 20:1 ratio if the blood pH is within the normal range. With 20 times more bicarbonate than carbonic acid, this capture system is most efficient at buffering changes that would make the blood more acidic. This is useful because most of the body’s metabolic wastes, such as lactic acid and ketones, are acids. Carbonic acid levels in the blood are controlled by the expiration of CO2 through the lungs. In red blood cells, carbonic anhydrase forces the dissociation of the acid into CO2 + H2O. CO2 diffuses into the alveoli, which shifts the system to the left and lowers H+, rendering the blood less acidic (Figure 14.3.2).
this figure shows has exchange between an alveolus and blood in the pulmonary capillaries. the directions of diffusion of oxygen and carbon dioxide are shown, along with the equation for the carbonic acid-bicarbonate buffer system
Figure 14.3.2 – Gas Exchange Between the Alveoli and the Blood: At the level of the lungs, the carbonic acid-bicarbonate buffer system shifts to the left as carbon dioxide diffuses into the alveoli.

The level of bicarbonate in the blood is controlled via the kidneys, where bicarbonate ions in the renal filtrate are reabsorbed and passed back into the blood. However, the bicarbonate buffer is the primary buffering system of the IF surrounding the cells in tissues throughout the body.

Respiratory Regulation of Acid-Base Balance

One function of the respiratory system is to eliminate the CO2 produced during cellular respiration. In healthy individuals, CO2 is formed at the same rate at which it is expelled, leaving H2CO3 unchanged:

CO2 + H2O ↔ H2CO3 ↔ H+ + HCO3−

 

The respiratory system can alter plasma pH rapidly (seconds to minutes) by altering ventilation rate and changing plasma PCO2 (Figure 14.3.3); this in turn will change plasma H+ by changing the rate of the forward or reverse reaction in the carbonic acid-bicarbonate buffer system:
  • hyperventilation -> decreased plasma CO2 -> decreased plasma H+ -> increased plasma pH
  • hypoventilation -> increased plasma CO2 -> increased plasma H+ -> decreased plasma pH

Note that hyperventilation and hypoventilation can be the causes acid-base imbalances, or they can be compensations to correct acid base imbalances (Figure 14.3.3.) A compensation when referring to acid-base balance is a physiological response that corrects and acid-base imbalance.

This top to bottom flowchart describes the regulation of PH in the blood. The left branch shows acidosis, which is when the PH of the blood drops. Acidosis stimulates brain and arterial receptors, triggering an increase in respiratory rate. This causes a drop in blood CO two and H two CO three. A drop in these two acidic compounds causes the blood PH to rise back to homeostatic levels. The right branch shows alkalosis which is when the PH of the blood rises. Alkalosis also stimulates brain and arterial receptors, but these now trigger a decrease in respiratory rate. This causes an increase in blood CO two and H two CO three, which lowers the PH of the blood back to homeostatic levels.
Figure 14.3.3 – Respiratory Regulation of Blood pH: The respiratory system can alter blood pH by changing ventilation rate.

Renal Regulation of Acid-Base Balance

The renal capacity to maintain acid-base balance is large but slow (minutes to hours); renal regulation of the body’s acid-base balance addresses the metabolic component of the buffering system. While the respiratory system (together with the respiratory center in the medulla oblongata) controls the blood levels of carbonic acid by controlling the exhalation of CO2, the kidneys can alter blood HCO3− regulate blood pH. The PCT, DCT, and collecting duct are the key areas of the nephron important in acid-base balance. Complex transport mechanisms in these areas regulate levels of HCO3− and H+. In this text, we focus on the activities in the PCT that are important to acid-base balance.

The cells of the PCT reabsorb HCO3– and actively secrete H+ into the filtrate. Bicarbonate ions found in the filtrate are essential to the bicarbonate buffer system, yet the cells of the tubule are not permeable to bicarbonate ions. The steps involved in supplying bicarbonate ions to the system are seen in Figure 14.3.4 and are summarized below:

  • Step 1: In the lumen of the PCT, HCO3– combines with H+ to form carbonic acid.
    • Note the source of H+: Na+ ions are reabsorbed from the filtrate in exchange for H+ by an antiport mechanism in the apical membranes of cells lining the renal tubule.
    • Carbonic anhydrase catalyzes the dissociation of carbonic acid into CO2 and water, which diffuse across the apical membrane into the cell.
  • Step 2: Inside the cell, the reverse reaction occurs to produce HCO3–
    • These bicarbonate ions are cotransported with Na+ across the basolateral membrane to the interstitial space around the PCT, and HCO3– moves into the blood (= HCO3– reabsorption.)
    • The cotransport of HCO3– with Na+ on the basolateral side is not shown in Figure 14.3.4
This diagram depicts a cross section of the left wall of a kidney proximal tubule. The wall is composed of two block-shaped cells arranged vertically one on top of each other. The lumen of the proximal tubule is to the left of the two cells. Yellow-colored urine is flowing through the lumen. There is a small strip of blue interstitial fluid to the right of the two cells. To the right of the interstitial fluid is a cross section of a blood vessel. A loop of chemical reactions is occurring in the diagram. Within the lumen of the proximal tubule, HCO three minus is combining with an H plus ion that enters the lumen from a proximal tubule cell. This reaction forms H two CO three. H two CO three then breaks into H two O and CO two, a reaction catalyzed by the enzyme carbonic anhydrase. The CO two then moves from the lumen of the proximal tubule into one of the proximal tubule cells. There, the reaction runs in reverse, with CO two combining with H two O to form H two CO three. The H two CO three then splits into H plus and HCO three minus. The H plus moves into the lumen, reinitiating the first step of the loop. The HCO three minus leaves the proximal tubule cell and enters the blood stream.
Figure 14.3.4 – Reabsorption of HCO3– and Secretion of H+ by the Cells of the PCT: Tubular cells are not permeable to bicarbonate; thus, bicarbonate is conserved rather than reabsorbed. This figure is a simplification of the process occurring in the PCT.

With certain acid-base imbalances, renal mechanisms will be the main type of compensations used to correct the imbalance:

  • if plasma pH is too low: the kidneys increase HCO3– reabsorption and increase H+ secretion, which results in an increase in plasma pH
  • if plasma pH is too high: the kidneys decrease HCO3– reabsorption and decrease H+ secretion, which results in a decrease in plasma pH

Section Review

A variety of buffering systems exist in the body that helps maintain the pH of the plasma within a narrow range—between pH 7.35 and 7.45. A buffer is a substance that prevents drastic changes in fluid pH by absorbing excess hydrogen or hydroxyl ions. Several substances serve as buffers in the body, including cell and plasma proteins, hemoglobin, phosphates, bicarbonate ions, and carbonic acid.

The carbonic acid-bicarbonate buffer system is the primary buffering system of the ECF, and proteins are the primary buffering system in the ICF. The respiratory and renal systems play major roles in acid-base homeostasis.

Review Questions

Critical Thinking Questions

Glossary

buffer
a substance or chemical system that prevents a drastic change in pH of a fluid
compensation
a physiological response that corrects an acid-base imbalance

Glossary Flashcards

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