8.10: Compensation of pH via Organs

Learning Objectives

• Describe the role of lungs and the kidneys in maintaining a stable blood serum pH with the buffers

The pH of your blood normally ranges between 7.35 and 7.45. A blood pH below the normal range is called acidosis. A blood pH above this range is called alkalosis. Either one is potentially fatal. Why? Enzymes catalyze chemical reactions in the body. The enzymes are made of proteins which are chains of amino acids connected via peptide bonds. Some of the amino acids in the proteins contain acidic or basic functional groups. Changes in the pH of the blood serum can change the nature of the functional groups in the side chain through neutralization reactions. This in turn will change the intermolecular forces that affect protein folding or shape of the protein. The change in the shape of the protein will adversely affect the function of the enzyme or protein.

To maintain a constant serum pH the body relies on two approaches.

• One involves the use of buffers to minimize the changes in hydronium ion concentrations. This was discussed in the earlier section 8.9.
• The other is to control the hydronium ion concentrations by the involvement of the organs such as lungs and the kidneys.

Carbonic Acid and Bicarbonate Ion Buffer System

The most important buffer system in blood is formed from carbonic acid (H₂CO₃) and its conjugate base, the bicarbonate ion (HCO₃^-). Carbon dioxide (CO₂) is a product of metabolism. Some of the CO₂ dissolves in the blood serum to generate carbonic acid as shown in equation 1. The carbonic acid then acts as a Bronsted-Lowry acid to set up an equilibrium as shown in equation 2.

Equation 1:                     H₂O(l)  +   CO₂(g)    ⇌     H₂CO₃(aq)

Equation 2:                     H₂CO₃(aq)   +    H₂O(l)     ⇌      HCO₃^-(aq)   +    H₃O⁺(aq)

Acidosis

During conditions of acidosis, when the pH falls below 7.35, the concentration of H₃O⁺ increases, the equilibrium in equation 2 shifts left increasing the concentration of H₂CO₃. Equilibrium in equation 1 then shifts left as a result of increasing H₂CO₃ concentrations, increasing the concentration of CO₂. The CO₂ is then expelled out by increased respiration to lower CO₂. This is how the lungs compensate for acidosis.

The kidney play, a role during conditions of acidosis, when the pH falls below 7.35. As the concentration of H₃O⁺ increases, the equilibrium in the equation shifts left, decreasing the concentration of HCO₃^- ion. The kidneys can then remove the excess H₃O⁺ into the urine or add HCO₃^- ion into the blood stream, thus correcting the pH. This is how the kidneys compensate for acidosis.

The role of the organs to compensate for acidosis are summarized in figure 8.10.1 below.

Figure \(\PageIndex{1}\): Role of organs to compensate for acidosis.

Alkalosis

During conditions of alkalosis, when the pH increases above 7.45, the relative concentration of H₃O⁺ decreases, the equilibrium in equation 2 shifts right decreasing the concentration of H₂CO₃. When the concentration of H₂CO₃ decreases the equilibrium in equation 1 shifts right decreasing the concentration of CO₂. The respiration is then decreased to conserve CO₂. This is compensation for alkalosis by the lungs.

Equation 1:                     H₂O(l)  +   CO₂(g)    ⇌     H₂CO₃(aq)

Equation 2:                     H₂CO₃(aq)   +    H₂O(l)     ⇌      HCO₃^-(aq)   +    H₃O⁺(aq)

The kidney play, a role during conditions of alkalosis, when the pH is above 7.35. As the concentration of H₃O⁺ decreases, the equilibrium in equation 2 shifts right increasing the concentration of HCO₃^- ion. The kidneys can then add H₃O⁺ into the blood stream and remove excess the HCO₃^- ion through the urine, correcting the pH. This is compensation for alkalosis by the kidneys.

The role of the organs to compensate for alkalosis are summarized in figure 8.10.2 below.

Figure \(\PageIndex{2}\): Role of organs to compensate for alkalosis.

Dihydrogenphosphate and the Hydrogenphosphate Buffer System

The second buffer system is the dihydrogenphosphate (H₂PO₄^-) and its conjugate base the (HPO₄^2-). The mixture of anions H₂PO₄^- / HPO₄^2- sets up the following acid-base equilibrium. Note the absence of gases in this equilibrium and the presence of three ions in the equilibrium.

H₂PO₄^-(aq)  +   H₂O(l)    ⇌     HPO₄^2-(aq)  +   H₃O⁺(aq)

Acidosis

During conditions of acidosis, when the pH falls below 7.35, the concentration of H₃O⁺ increases, the equilibrium in equation shifts left increasing the concentration of H₂PO₄^- and decreasing the concentrations of HPO₄^2-. The kidneys remove excess H₃O⁺ and H₂PO₄^- via urine and add HPO₄^2- to the blood. This is another compensation for acidosis by the kidneys using a second buffer system.

Alkalosis

During conditions of alkalosis, when the pH is above 7.45, the relative concentration of H₃O⁺ decreases, the equilibrium in the equation shifts right increasing the concentration of HPO₄^2- ion and decreasing the concentrations of H₂PO₄^- ion. The kidneys can then add H₃O⁺ and H₂PO₄^- into the blood stream and remove excess the HPO₄^2- ion through the urine, correcting the pH. This is compensation for alkalosis by the kidneys.

The role of the kidneys to compensate for acidosis and alkalosis are summarized in the diagram below.

Note

Even with all these controls, the pH of blood can still sometimes fall outside the normal range. The acid/base disorders are classified as being either respiratory or metabolic in nature. Respiratory disorders occur when exhaling does not remove CO₂ from the blood at the same rate that it is produced in cells. Metabolic disorders arise from the inability to remove acids produced by metabolism, inability to control concentrations of HCO₃^-, ingestion of compounds that are acids or bases or compounds that can be converted to acids or bases. The following lists conditions that can lead to either acidosis or alkalosis.

Acidosis:

• Lung disease, asthma
• Excessive alcohol consumption
• Ketosis, ketoacidosis (starvation, diabetic conditions)
• Holding your breath too long
• Hypoventilation (not enough CO₂ is exhaled)

Alkalosis:

• Excessive use of antacids
• Anxiety, fever
• Hyperventilation (CO₂ is blown off faster than it is produced in the cells)

Mild cases of acidosis can lead to light-headedness. Severe cases can lead to coma and death. Symptoms of alkalosis include headaches, nervousness, cramps. In severe cases of alkalosis convulsions and death results.