Question

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

Answer 1

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?