The pH scale is an artificial scale used to distinguish between acid, neutral, and basic/alkaline solutions.

  • The pH scale is a convenient measure of the $\rm [H^+]$ in a solution, which enables different solutions to be compared.
  • $\rm pH = –\log_{10} [H^+]$; $\rm [H^+] = 10^{–pH}$
  • At $\rm 25°C/298~K$:
    $\rm pH < 7: [H^+] > [OH^–] = acidic$
    $\rm pH = 7: [H^+] = [OH^–] = neutral$
    $\rm pH > 7: [H^+] < [OH^–] = basic$
  • An increase of one pH unit corresponds to a ten-fold decrease in $\rm [H^+]$.
    A decrease of one $\rm pH$ unit corresponds to a ten-fold increase in $\rm [H^+]$.
  • $\rm pH$ can be determined using universal indicator and pH meters. $\rm pH$ meters are more precise.
  • Water is a weak electrolyte. The dissociation of water can be represented by the equilibrium:
    $\rm H_2O(l) \rightleftharpoons H^+(aq) + OH^–(aq)$
    with the equilibrium constant: $\mathrm K_c = \rm \dfrac{[H^+(aq)][OH^–(aq)]}{[H_2O(l)]}$
    This can be simplified to $\mathrm K_w = \mathrm K_c\rm [H_2O(l)] = [H^+(aq)][OH^–(aq)]$
    As $\mathrm K_w$, is an equilibrium constant it depends on the temperature.
  • The ionic product constant of water, $\mathrm K_w$, has a fixed value at a fixed temperature.
    $\mathrm K_w \rm = [H^+] [OH^–]$
  • At $\rm 25°C$: $\mathrm K_w = \rm [H^+] [OH^–] = 10^{–14}$ at $\rm 298~K$
    For a neutral solution: $\rm [H^+] = [OH^–] = 10^{–7}$  
    $\rm pH = –\log_{10} 10^{–7} = 7$
  • As $\rm [H^+]$ increases $\rm [OH^–]$ decreases in aqueous solution.
    As $\rm [OH^–]$ increases $\rm [H^+]$ decreases in aqueous solution.