How is acid/base chemistry defined and represented in aqueous solution?

1 Answer
Mar 23, 2017

Answer:

In water under standard conditions, we know that #K_a=[H_3O^+][HO^-]=10^-14#

(note that #H_3O^+# and #H^+# may be used equivalently. They represent the characteristic cation in water, #"the acidium ion."#)

Explanation:

We know (i) that #pH=-log_10[H_3O^+]#, and (ii) that under standard conditions of #1*atm# and #298*K#, #K_a=10^-14#.

(i) was by definition, and (ii) was by dint of VERY careful measurement. And these relationships define the very idea of acid/base chemistry in aqueous solution.

A solution is acidic if there there is a preponderance of the characteristic cation, #H_3O^+#. That is the solution is acidic if #[H_3O^+]>[HO^-]#, i.e. there are more so-called #"hydronium ions"# than #"hydroxide ions"# per unit volume.

The solution is basic if #[H_3O^+]<[HO^-]#, i.e. there are LESS so-called #"hydronium ions"# than #"hydroxide ions"# per units volume.

And if #[H_3O^+]=[HO^-]#, then the solution is NEUTRAL.

#"For undergrads:"#

we know that #K_a=[H_3O^+][HO^-]=10^-14# under standard conditions of #1*atm# (or whatever they use these days), and #298*K#. Under non-standard conditions, i.e. #T=373*K#, how do you think #K_a# would evolve?