he pH of a Solution Is Determined by the Relative Concentrations of Acids and Bases. The relationship between the pH of a solution and the concentrations of an acid and its conjugate base is easily derived. Equation 2-5 can be rearranged to [2-7] Taking the negative log of each term (and letting pH ⫽⫺log[H ⫹ ]; Eq. 2-3) gives [2-8] Substituting pK for ⫺log K (Eq. 2-6) yields [2-9] This relationship is known as the Henderson–Hasselbalch equation. When the molar concentrations of an acid (HA) and its conjugate base (A ⫺ ) are equal, log ([A ⫺ ]兾[HA]) ⫽ log 1 ⫽ 0, and the pH of the solution is numerically equiv- alent to the pK of the acid. The Henderson–Hasselbalch equation is invaluable for calculating, for example, the pH of a solution containing known quanti- ties of a weak acid and its conjugate base (see Sample Calculation 2-2). However, since the Henderson–Hasselbalch equation does not account for the ionization of water itself, it is not useful for calculating the pH of solu- tions of strong acids or bases. For example, in a 1 M solution of a strong acid, [H ⫹ ] ⫽ 1 M and the pH is 0. In a 1 M solution of a strong base, [OH ⫺ ] ⫽ 1 M, so [H ⫹ ] ⫽ [OH ⫺ ]兾K w ⫽ 1 ⫻ 10 ⫺14 M and the pH is 14.