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Water naturally dissociates into an equilibrium mixture of $H^+$ and $OH^-$ ions and $H_2O$ molecules
$$
H_2O \rightleftharpoons^{K_W} H^++OH^-\,,
$$
where the concentrations of $H^+$ and $OH^-$ ions, written as $[H^+]$ and $[OH^-]$ are related by the expression
$$
K_W = [H^+][OH^-].
$$
$K_W$ is called the dissociation constant, and depends on the temperature of the water.
The following table of data shows the dissociation constant for water at various temperatures and standard pressure.
| Water temperature | $\quad K_W\times10^{14}\quad$ |
| $0^\circ$ C | 0.1 |
| $10^\circ$ C | 0.3 |
| $18^\circ$ C | 0.7 |
| $25^\circ$ C | 1.2 |
| $30^\circ$ C | 1.8 |
| $50^\circ$ C | 8.0 |
| $60^\circ$ C | 13 |
| $70^\circ$ C | 21 |
| $80^\circ$ C | 35 |
| $90^\circ$ C | 53 |
| $100^\circ$ C | 73 |
From this table, work out an estimate for the temperature at which water has a $pH$ of exactly 7, 6.8 and 7.2. Recall that the $pH$ is defined as $pH=-\log_{10}([H^+])$