Biochemical reactions are specifically sensitive to pH. Most biological molecules contain groups of atoms that may be charged or neutral depending on pH, and whether these groups are charged or neutral has a significant effect on the biological activity of the molecule. In all multicellular organisms, the fluid within the cell and the fluids surrounding the cells have a characteristic and nearly constant pH. There is great variation in the pH of fluids in the body and small variation is found within each system. For example, the pH of body fluid can vary from 8 in the pancreatic fluid to 1 in the stomach. The average pH of blood is 7.4, and of cells is in the range of 7.3 to 7.
This pH of body fluids is maintained through buffer systems. Body fluids contain buffering agents and buffer systems that maintain pH at or near 7.4. 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. Proteins are the most important buffers in the body as their amino and carboxylic acid groups act as proton donors or acceptors as H+ ions are either added or taken out from the environment. Important endogenous (natural) buffer systems include carbonic acid/sodium bicarbonate and sodium phosphate in the plasma and hemoglobin, and potassium phosphate in the cells.
Two important biological buffer systems are the dihydrogen phosphate system and the carbonic acid system.
1. The Phosphate Buffer System:
The phosphate buffer system operates in the internal fluid of all cells. This buffer system consists of dihydrogen phosphate ions (H2PO–4 ) as hydrogen ion donor (acid) and hydrogen phosphate ions (HPO2-4) as hydrogen-ion acceptor (base). These two ions are in equilibrium with each other as indicated by the chemical equation given below.
If additional hydrogen ions enter the cellular fluid, they are consumed in the reaction with H2PO–4, and the equilibrium shifts to the left. If additional hydroxide ions enter the cellular fluid, they react with H2PO–4, producing HPO2-4, and shifting the equilibrium to the right. The equilibrium expression for this equilibrium is expressed as given below.
The value of Ka for this equilibrium is 6.23 × 10-8 at 25°C. From equation (1), the relationship between the hydrogen ion concentration and the concentrations of the acid and base can be derived as follows.
Thus, when the concentrations of H2PO–4 and HPO2-4 are the same, the value of the molar concentration of hydrogen ions is equal to the value of the equilibrium constant, and therefore;
pH = pKa (− log Ka) …equation (3)
Buffer solutions are most effective in maintaining a pH near the value of the pKa. In mammals, cellular fluid has a pH in the range of 6.9 to 7.4, and the phosphate buffer is effective in maintaining this pH range. The pKa 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 because H3PO4 and H2PO–4 are found in very low concentrations in the blood. Hemoglobin also acts as a pH buffer in the blood. Hemoglobin protein can reversibly bind either H+ (to the protein) or O2 (to the Fe of the heme group), but that when one of these substances is bound, the other is released. 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.
2. The Carbonic Acid System:
Another biological fluid in which a buffer plays an important role in maintaining pH in blood plasma. In blood plasma, the carbonic acid and hydrogen carbonate ion equilibrium buffers the pH. In this buffer, carbonic acid (H2CO3) is the hydrogen ion donor (acid) and hydrogen carbonate ion (HCO–3 ) is the hydrogen-ion acceptor (base). The simultaneous equilibrium reaction is shown below.
This buffer functions in the same way as the phosphate buffer. Additional H+ is consumed by HCO–3 and additional OH– is consumed by H2CO3. The value of Ka for this equilibrium is 7.9 × 10-7, and the pKa is 6.1 at body temperature. In blood plasma, the concentration of hydrogen carbonate ion is about twenty times the concentration of carbonic acid. The pH of arterial blood plasma is 7.4. If pH falls below this normal value, a condition called acidosis and when pH rises above the normal value, the condition is called alkalosis is observed.
The concentrations of hydrogen carbonate ions and carbonic acid are controlled by two physiological systems. The concentration of hydrogen carbonate ions is controlled through the kidneys whereas excess hydrogen carbonate ions are excreted in the urine. The carbonic acid-hydrogen carbonate ion buffer works throughout the body to maintain the pH of blood plasma close to 7.4. Changes in hydrogen carbonate ion concentration, however, require hours through the relatively slow elimination through the kidneys. Carbonic acid concentration is controlled by respiration that is through the lungs. Carbonic acid is in equilibrium with dissolved carbon dioxide gas. An enzyme called carbonic anhydrase catalyzes the conversion of carbonic acid to dissolved carbon dioxide. In the lungs, excess dissolved carbon dioxide is exhaled as carbon dioxide gas.
The body maintains the buffer by eliminating either the acid (carbonic acid) or the base (hydrogen carbonate ions). Changes in carbonic acid concentration bring about within seconds through increased or decreased respiration.
A lysis buffer is used for lysing cells for use in experiments that analyze the compounds of the cells (for example, western blot). There are many kinds of lysis buffers that one can apply; depending on what analysis the cell lysate will be used, for example, RBC lysis buffer. In studies like DNA fingerprinting, the lysis buffer is used for DNA isolation. Dish soap can be used in a pinch to break down the cell and nuclear membranes, allowing the DNA to be released.
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