Chemicals in the Body
All cells in the body continually exchange chemicals (e.g.,nutrients, waste products, and ions) with the external fluid surrounding them. This external fluid, in turn, exchanges chemicals with the blood being pumped throughout the body. Hence, the chemical composition of the blood (and therefore of the external fluid) is extremelyimportant for the cell. If, for instance, the pH of the blood and external fluid is too low (too many H+ ions), then an excess of H+ ions will enter the cell. Maintaining the proper pH is critical for the chemical reactions that occur in the body. In order to maintain the proper chemical composition inside the cells, the chemical composition of the fluids outside the cells must be kept relativelyconstant. The most important way that the pH of the blood is kept relatively constant is by buffers dissolved in the blood.
How Buffers Work
Acid-base buffers confer resistance to a change in the pH of a solution when hydrogen ions or hydroxide ions are added or removed. An acid-base buffer typically consists of a weak acid, and its conjugate base (salt). Buffers work because the concentrationsof the weak acid and its salt are large compared to the amount of protons or hydroxide ions added or removed. When protons are added to the solution from an external source, some of the base component of the buffer is converted to the weak-acid component; when hydroxide ions are added to the solution protons are dissociated from some of the weak-acid molecules of the buffer, converting them tothe base of the buffer. However, the change in acid and base concentrations is small relative to the amounts of these species present in solution. Hence, the ratio of acid to base changes only slightly.
By far the most important buffer for maintaining acid-base balance in the blood is the carbonic-acid-bicarbonate buffer. The simultaneous equilibrium reactions of interest are:
Carbonic acid(H2CO3) is the acid and water is the base. The conjugate base for H2CO3 is HCO3..Carbonic acid also dissociates rapidly to produce water and carbon dioxide. This second process is not an acid-base reaction, but it is important to the blood's buffering capacity.
In the equation above, pK is equal to the negative log of the equilibrium constant, K, for the buffer.
where K=Ka/K2 .
Thisquantity provides an indication of the degree to which HCO3- reacts with H+ to form H2CO3, and subsequently to form CO2 and H2O. In the case of the carbonic-acid-bicarbonate buffer, pK=6.1 at normal body temperature.
Derivation of the pH Equation for the Carbonic-Acid-Bicarbonate Buffer
We may begin by defining the equilibrium constant, K1, using the Law of Mass Action:
Ka is the equilibrium constantfor the acid-base reaction. The formula for Ka is:
The equilibrium constant, K2 is also defined by the Law of Mass Action:
Because the two equilibrium reactions occur simultaneously, the 2 equations above can be treated as two simultaneous equations.
Rearranging that equation allows us to solve for the equilibrium proton concentration in terms of the two equilibrium constants and theconcentrations of the other species:
Because we are interested in the pH of the blood, we take the negative log of both sides of the equation.
Recalling the definitions of pH and pK, the last equation can be rewritten:
The pH of the buffered solution is dependent only on the ratio of the amount of CO2 present in the blood to the amount of HCO3- (bicarbonate ion) present in the blood. This ratioremains relatively constant, because the concentrations of both buffer components (HCO3- and CO2) are very large, compared to the amount of H+ added to the blood during normal activities and moderate exercise. When H+ is added to the blood as a result of metabolic processes, the amount of HCO3- decreases; however, the amount of the change is tiny compared to the amount of HCO3- present in the blood....