Acid-base chemistry is subtle. The dissociation of water leads to formation of charged species known as hydronium and hydroxide ion, H3O+ and OH-, respectively. However, the equilibrium concentrations are so small that you might think at first that this equilibrium is unimportant. The concentration of each ion in neutral water is 10-7 M. Why are such low concentrations important? Protonation and deprotonation leads to catalysis. The balance of the concentration of these ions determines reactivity and the balance of other ions that also lead to chemical reactions in water and mixed solvent. Even species with a relatively low concentratoin can play an important role under these circumstances. The dynamic range of the H3P+ concentration is quite large. It can vary from 10-14 to 1 M. The same is true for hydroxide. Since the reactivity of these and other ions are modulated by these concnetrations we need a way to represent the concnetration of this large range. Hence, we adopt a logarithmic scale, the p-scale. We define the pX of some concentration or equilibrium constant as follows:
pX = -log10 X
These definitions and the general terms used in acid-base chemistry are described in the following segment.
It is interesting to note that we have some experience with the large dynamic range of hydronium ion concentration in our everyday life. Some observations are summarized in the following segment that describes commonly observed substances or processes that depend on solution pH ranging from 1-14.
