(also buffer solutions or buffer mixtures), systems that maintain a certain concentration of H+ ions—that is, a certain acidity of the medium. The acidity of buffer solutions changes little when they are diluted or when certain amounts of acids or bases are added.
A mixture of solutions of acetic acid CH3COOH and its sodium salt CH3COONa is an example of a buffer system. This salt, a strong electrolyte, dissociates almost entirely—that is, it yields many CH3COO− ions. When a strong acid that yields many H+ ions is added to a buffer system, the ions are bound to CH3COO− ions and form weak (that is, only slightly dissociating) acetic acid:
(1) H1 + CH3COO- ⇆ CH3COOH
On the other hand, when a buffer system is alkalized—that is, when a strong base (for example, NaOH) is added, the OH− ions are bound to the H+ ions present in the buffer system owing to the dissociation of acetic acid and form a weak electrolyte—water:
(2) H+OH ⇆ H2O
As the H+ ions are expended in binding the OH− ions, newer and newer CH3COOH molecules dissociate so that the equilibrium (1) shifts to the left. As a result, both when H+ and OH− ions are added, these ions are bound, and consequently, the acidity of the solution scarcely changes.
The acidity of solutions is usually expressed by the so-called hydrogen ion exponent or pH (pH = 7 for neutral solutions, below 7 for acid solutions, and above 7 for alkaline solutions). Adding 100 ml of an 0.01 molar solution of HCl (0.01 M) to 1 liter of pure water lowers the pH from 7 to 3. Adding the same solution to 1 liter of the buffer system CH3COOH CH3COONa (0.1 M) lowers the pH from 4.7 to 4.65—that is, by only 0.05. In the presence of 100 ml of 0.01 M solution of NaOH in pure water, the pH changes from 7 to 11, but only from 4.7 to 4.8 in the above buffer system. There are many other buffer systems in addition to those discussed above (see Table 1). The acidity (and consequently the pH) of a buffer system varies with the nature of the components and their concentration and, in some systems, with the temperature as well. The pH of every buffer system remains more or less constant only to a certain limit, which varies with the concentration of the constituents.
|Table 1. Examples of buffer systems|
|Constituents (concentrations of 0.1 g-mole/l)||pH (at 15°-25°C)|
|Acetic acid + sodium acetate, CH3COOH + CH3COONa...............||4.7|
|Sodium citrate (disubstituted), C6H6O7Na2...............||5.0|
|Boric acid + borax, H3BO3 + Na2B4O7 · 10H2O...............||8.5|
|Boric acid + sodium hydroxide, H3BO3 + NaOH...............||9.2|
|Sodium phosphate (disubstituted) + sodium hydroxide, Na2HPO4 + NaOH...............||11.5|
Buffer systems are widely used in analytical practice and in chemical production because many chemical reactions proceed in the desired direction and at the appropriate speed only within narrow pH limits. Buffer systems are important in the life processes of organisms. They keep the acidity of various biological fluids constant (blood, lymph, intercellular fluids). The main buffer systems of animals and man are the bicarbonate (carbonic acid and its salts) and phosphate (phosphoric acid and its salts) systems and proteins (their buffer properties are determined by the presence of basic and acidic groups). Blood proteins (chiefly hemoglobin, which accounts for about 75 percent of the buffer capacity of the blood) keep the blood pH fairly stable. The blood pH in man varies from 7.35 to 7.47 and remains in this range even after substantial changes in nutrition and other conditions. To shift the blood pH to the alkaline side requires the addition to it of 40-70 times more alkali than to an equal volume of pure water. Natural buffer systems in the soil play a major role in maintaining its fertility.
V. L. VASILEVSKII