Buffering against pH Changes in Biological Systems: -Weak Acids or Bases Buffer Cells and Tissues against pH Changes

The intracellular and extracellular fluids of multicellular organisms have a characteristic and nearly constant pH. The organism’s first line of defense against changes in internal pH is provided by buffer systems. The cytoplasm of most cells contains high concentrations of proteins, which contain many amino acids with functional groups that are weak acids or weak bases. For example, the side chain of histidine (Fig. 2–20) has a pK_a of 6.0; proteins containing histidine residues therefore buffer effectively near neutral pH. Nucleotides such as ATP, as well as many low molecular weight metabolites, contain ionizable groups that can contribute buffering power to the cytoplasm. Some highly specialized organelles and extracellular compartments have high concentrations of compounds that contribute buffering capacity: organic acids buffer the vacuoles of plant cells; ammonia buffers urine.

FIGURE 2–20 The amino acid histidine, a component of proteins, is a weak acid. The pK_a of the protonated nitrogen of the side chain is 6.0.

Two especially important biological buffers are the phosphate and bicarbonate systems. The phosphate buffer system, which acts in the cytoplasm of all cells, consists of H₂PO₄^- as proton donor and HPO₄^-2 as pro ton acceptor:

The phosphate buffer system is maximally effective at a pH close to its pKa of 6.86 (Figs 2–16, 2–18) and thus tends to resist pH changes in the range between about 5.9 and 7.9. It is therefore an effective buffer in biological fluids; in mammals, for example, extracellular flu ids and most cytoplasmic compartments have a pH in the range of 6.9 to 7.4. Blood plasma is buffered in part by the bicarbonate system, consisting of carbonic acid (H₂CO₃) as proton donor and bicarbonate (HCO₃) as proton acceptor:

This buffer system is more complex than other conjugate acid-base pairs because one of its components, carbonic acid (H₂CO₃), is formed from dissolved (d) car bon dioxide and water, in a reversible reaction:

Carbon dioxide is a gas under normal conditions, and the concentration of dissolved CO₂ is the result of equilibration with CO₂ of the gas (g) phase:

The pH of a bicarbonate buffer system depends on the concentration of H₂CO₃ and HCO₃, the proton donor and acceptor components. The concentration of H2CO3 in turn depends on the concentration of dissolved CO₂, which in turn depends on the concentration of CO₂ in the gas phase, called the partial pressure of CO₂. Thus the pH of a bicarbonate buffer exposed to a gas phase is ultimately determined by the concentration of HCO₃ in the aqueous phase and the partial pressure of CO₂ in the gas phase (Box 2–4). Human blood plasma normally has a pH close to 7.4. Should the pH-regulating mechanisms fail or be over whelmed, as may happen in severe uncontrolled diabetes when an overproduction of metabolic acids causes acidosis, the pH of the blood can fall to 6.8 or below, leading to irreparable cell damage and death. In other diseases the pH may rise to lethal levels.

FIGURE 2–21 The pH optima of some enzymes. Pepsin is a digestive enzyme secreted into gastric juice; trypsin, a digestive enzyme that acts in the small intestine; alkaline phosphatase of bone tissue, a hydrolytic enzyme thought to aid in bone mineralization.

Although many aspects of cell structure and function are influenced by pH, it is the catalytic activity of enzymes that is especially sensitive. Enzymes typically show maximal catalytic activity at a characteristic pH, called the pH optimum (Fig. 2–21). On either side of the optimum pH their catalytic activity often declines sharply. Thus, a small change in pH can make a large difference in the rate of some crucial enzyme-catalyzed reactions. Biological control of the pH of cells and body fluids is therefore of central importance in all aspects of metabolism and cellular activities.

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