The kidneys and the lungs work together to help maintain a blood pH of 7.4 by affecting the components of the buffers in the blood. To understand how these organs help control the pH of the blood, we must first discuss how buffers work in solution.

Acid-base buffers confer resistance to a change in the pH of a solution when hydrogen ions (protons) or hydroxide ions are added or removed. An acid-base buffer consists of a weak acid and its conjugate base (from a salt) or a weak base and its conjugate acid (from a salt). A buffer works because it contains a substantial amount of a weak acid and a weak base (the conjugate acid-base pair) at equilibrium with each other. When protons (from an external source) are added to the buffer, some of the base component of the buffer will react with the protons and turn into the conjugate acid (which is the weak-acid component of the buffer) and thus neutralizing most of the protons added. When hydroxide ions are added (or, equivalently, when protons are removed from the buffer), some of the weak-acid component of the buffer will dissociate and turn into the conjugate base (which is the weak-base component of the buffer) thus replenishing most of the protons removed. Hence, adding a small amount of acid or base to a buffer solution merely changes the ratio of the conjugate acid and conjugate base in an acid-base equilibrium. Thus, the effect on the pH of the solution is small, within certain limitations of the amount of H+or OH-added or removed.




Hence, the physiological blood pH of 7.4 is outside the optimal buffering range of the carbonic acid-bicarbonate buffer. The addition of protons to the blood due to strenuous exercise may be too great for the buffer alone to effectively control the blood’s pH. When this happens, other organs must help control the amounts of CO2 and HCO 3 - in the blood. The lungs remove excess CO2 from the blood (helping to raise the pH as equilibria in Eq.10 shift to the left as predicted by the Le Chatelier’s Principle). When the pH of the body is excessively high (a condition known as alkalosis), the kidneys remove bicarbonate ion (HCO3-) from the blood (helping to lower the pH as equilibria in Eq. 10 shift to the right).

Other pH-Buffer Systems in the Blood

Other buffers perform minor roles than the carbonic-acid-bicarbonate buffer in regulating the pH of the blood. The phosphate buffer consists of H2PO4- in equilibrium with HPO42- and H+. The pK for the phosphate buffer is 6.8, which allows this buffer to function within its optimal buffering range at physiological pH. The phosphate buffer only plays a minor role in the blood, however, because H2PO4- and HPO42- are found in very low concentration in the blood. Hemoglobin also acts as a pH buffer in the blood. Recall from the “Hemoglobin and the Heme Group: Metal Complexes in the Blood for Oxygen Transport” tutorial from Chem 151 that the hemoglobin protein can reversibly bind either H+ (to the protein) or O2 (to the Fe of the heme group), and that when one of these substances is bound, the other is released (as explained by the Bohr effect). During exercise, hemoglobin helps to control the pH of the blood by binding some of the excess protons that are generated in the muscles. At the same time, molecular oxygen is released for use by the muscles.

Buffers in the blood resist pH change



What are buffers?



How do buffers work?



Carbonic acid-bicarbonate buffer system:



- When the blood becomes too acidic (low pH), the excess H+ ions combine with bicarbonate ions to form carbonic acid, shifting the equilibrium to the left. This helps remove the excess H+ ions and raises the pH back to normal.
- When the blood becomes too alkaline (high pH), the equilibrium shifts to the right, releasing H+ ions from carbonic acid and increasing the concentration of H+ ions in the blood. This lowers the pH back to normal.

Importance of buffers in the blood:





Conclusion: