Active Reaction

ACTIVE REACTION, the reaction of a solution determined by the presence in it of free hydrogen and hydroxyl ions. The determination of the acidity or alkalinity of a solution by the method of titration (see Alkalimetry) consists in measuring the amount of acid or alkali of known concentration which is required to neutralize the solution under examination. However, a comparison of solutions of different acids (or alkalis) having the same titrational acidity (or alkalinity) clearly shows that this method does not give a satisfactory measure of the true reaction of the solution. Thus, for example, normal (that is, containing one gram-equivalent of H per liter) solutions of hydrochloric, sulfuric, tartaric, formic, and acetic acids are neutralized by the same amount of caustic soda, i.e., they have the same titrational acidity. Meanwhile, the acidic properties of the acids mentioned decrease rapidly in the order in which they are listed above: normal acetic acid produces approximately the same effect as a 250 times more dilute solution of hydrochloric acid. According to the theory of electrolytic dissociation, the properties of electrolytes depend on the ions (see) into which they dissociate in solution. The properties of acids depend on the common hydrogen ion to them all. In the same way, the properties of alkalis are conditioned by the hydroxyl ion which they detach. The acidic or alkaline reaction is determined by the concentration of the free hydrogen or hydroxyl ions present in the solution. Strong acids are practically completely dissociated. In the case of weaker acids only a part of the molecules is dissociated; the amount of free, active H-ions is correspondingly smaller. For acetic acid, for example, dissociation proceeds according to the following equation, in which the charges of the cations (positively charged ions) are denoted by dots, and the anions (negatively charged ions) by commas: CH₃.COOH $CH₃.CO' + H\ In solution a large number of undissociated molecules CH₃.COOH is in equilibrium with a small amount of acetate and hydrogen ions. The equilibrium is disturbed during titration, because instead of the H-ions which are bound by the added alkali, new ones are immediately liberated, formed by the dissociation of the remaining acid molecules. Therefore titration does not permit the measurement of the true concentration of free H- or OH-ions, from which the true reaction of the solution depends. For the measurement of the latter other methods have been proposed which do not disturb the dissociation of the acid or alkali during the determination itself. It was further established that the very ions H' and OH' in aqueous solutions are in equilibrium with each other. In acidic solutions, along with an excess of H", a negligible but strictly definite amount of OH' ions is contained, and exactly the same extremely small amounts of H-ions are contained in any alkaline aqueous solution. The concentration of either of these ions can be easily calculated from the concentration of the other, which is inversely proportional to it. Since their ratio determines the true reaction of the solution, it is sufficient to indicate one concentration - that of H- or OH-ions to characterize the latter. For a number of reasons it has been customary to use the concentration of hydrogen ions or, according to the proposal of Sorensen, its decimal logarithm taken with the opposite sign (this logarithm is a negative quantity). This quantity received the name of the index of concentration of hydrogen ions, briefly the hydrogen index, and is denoted by the sign pH or Ph. As the following table shows, each value of pH corresponds to a definite concentration of H' and OH' ions (for 18°). pH [H-] | [OH'] pH [H-] [OH'] 2 3 4 5 10-¹ 10-» 10-" 10-' 10-¹ 0,74.10-" 0,74.10-" 0,74.10-" 0,74.10-" 0,74.10-' 10-' 10-' 10-" 10-* 10-" 0,74.10-" 0,74.10-' 0,74.10-' 0,74.10-" 0,74.10-' Thus, in an aqueous solution at any reaction both ions of water are present; their quantitative ratio entirely determines the reaction. If the concentration of H-ions is greater than that of OH'-the reaction is acidic, with the reverse ratio it is alkaline; neutrality is determined as the equality of both concentrations. According to the table above, a solution is neutral at a pH approximately equal to 7; at a lower pH the reaction is acidic, at a higher one it is alkaline. Moreover, the value given characterizes neutrality exactly only at 22°. At other temperatures slightly different values of neutrality are found: at 0°-about 7.5, at 18°-7.07, at 38°-6.74. The concept of reaction established at the present time differs essentially from the old conception based on the application of titrimetry. Reaction is determined by the concentration of actually existing, active (that is, in a free state and actively acting in chemical and biological processes) H- and OH-ions. Therefore it often receives the name of actual reaction, or, better, A. r. On the other hand, titration, as mentioned above, measures, along with the actual, free H- or OH-ions, the amount of undissociated, potential, reserve H- and OH-ions, which can be liberated by shifting the ionic equilibrium. This reserve alkalinity (see), or acidity, has an indirect significance, conditioning a greater or lesser stability of A. r. with respect to various influences disturbing it. A. r. exerts a tremendous influence on a multitude of vital processes. It is sufficient, for example, for a slight change in A. r. of the blood to sharply change the vital activity of various tissues, the course of respiratory processes, the work of the heart, etc. It is not surprising, therefore, that such an important quantity for life is carefully preserved by the organism at a strictly constant level. In man the pH of the blood lies always between 7.3 and 7.5 (on average, approx. 7.36), which corresponds to a weakly alkaline reaction. In maintaining the constancy of A. r. of the blood a significant role is played by its reserve alkalinity, thanks to which when acidic products are introduced into the blood new OH-ions are immediately liberated in place of those bound by the acid. As a result, blood (and similar fluids) acquire not only in the organism but even in vitro a greater stability with respect to any disturbances of their normal reaction (see Buffer properties). Therefore with the introduction of the concept of A. r. the titrimetric measurement of reserve acidity and alkalinity has not completely lost its significance.