The titration of a STRONG base with a STRONG acid

The pH of an initial solution of strong base is observed to undergo a precipitous transition as the solution is "titrated" with a solution of strong acid.

Ingredients: sodium hydroxide, hydrochloric acid

Procedure: A complete recipe follows.

1. Prepare a 1.00M aqueous solution of sodium hydroxide.

2. Prepare a 6.00M aqueous solution of hydrochloric acid.

3. Titrate 200.0mL of sodium hydroxide solution with the hydrochloric acid solution.

4. Observe the change in pH.

Understanding: The titration of the strong base, sodium hydroxide, with the strong acid, hydrogen chloride, has at its heart the neutralization reaction

NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)

The essence of the reaction is

base + acid → salt + water

As the sodium hydroxide and hydrogen chloride are both strong electrolytes, the reaction can also be represented

Na+(aq) + OH-(aq) + H+(aq) + Cl-(aq) → Na+(aq) + Cl-(aq) + H2O(l)

The net ionic equation for the neutralization reaction of the strong base titrated by the strong acid can be written

OH-(aq) + H+(aq) → H2O(l)         K = 1/Kw

The regime of excess strong base

At the start of the titration, the solution is observed to be strongly basic, with a pH exceeding 12. As the titration proceeds, the pH of the solution initially diminishes quite slowly. Why slowly? Because any strong acid that is added to the solution is immediately neutralized to salt and water, and the excess strong base keeps the solution strongly basic. The dominant species in the solution is the hydroxide ion, and the solution is basic.

The equivalence point

The initial solution has a certain number of equivalents of strong base. The point at which the number of equivalents of acid added equals the number of equivalents of base initially in the solution is called the equivalence point. As the equivalence point is approached, a dramatic reduction in the pH is observed due to the depletion of excess base. At the equivalence point, there is no excess strong base or strong acid. The dominant species in the solution is water and the chemical equilibrium that determines the pH of the solution is the autoionization of water

H2O(l) → OH-(aq) + H+(aq)         Kw = aH+ aOH- = 1.00 x 10-14 @ 25C

The activities of hydrogen ion and hydroxide ion are aH+ = aOH- = 1.00 x 10-7 so that pH = -log10 (aH+) = 7.000, and the solution is neutral.

The regime of excess strong acid

As additional strong acid is added to the solution, the formation of excess strong acid causes the pH of the solution to plummet further to the deeply acidic range of pH less than 2. The dominant species in the solution is the excess hydrogen ion, and the pH of the solution is strongly acidic.

QUALitative and QUANTitative understanding of titrations

Question: In the titration of a strong acid, hydrochloric acid, with the strong base, sodium hydroxide, the chemistry is quite straightforward. The regimes of chemistry are captured by two principal types of equilibria - strong acid/base solutions and the autoionization of water.

Consider the titration of 50.00mL of a 0.2000M solution of hydrochloric acid with a 0.2000 M solution of sodium hydroxide. Take Kw = 1.00 x 10-14.

For the initial strong acid solution, region IV, and region V, define (1) the dominant species in the solution, (2) the chemical equilibrium (or excess species) that defines the pH in that region, and (3) the equilibrium constant expression corresponding to that reaction (if appropriate).

Compute the pH of the solution at the following volumes of strong base added in the titration: 0.00, 10.00, 25.00, 40.00, 49.00, 49.90, 49.99, 50.00, 50.01, and 55.00mL.

You can check your answers here.