Titrations in Nonaqueous Solvents
Thus far we have assumed that the acid and base are in an aqueous solution. In- deed, water is the most common solvent in acid–base titrimetry. When considering the utility of a titration, however, the solvent’s influence cannot be ignored.
The dissociation, or autoprotolysis constant for a solvent, SH, relates the con- centration of the protonated solvent, SH2+, to that of the deprotonated solvent, S–. For amphoteric solvents, which can act as both proton donors and proton accep- tors, the autoprotolysis reaction is
You should recognize that Kw is just the specific form of Ks for water. The pH of a solution is now seen to be a general statement about the relative abundance of pro- tonated solvent
Perhaps the most obvious limitation imposed by Ks is the change in pH during a titration. To see why this is so, let’s consider the titration of a 50 mL solution of 10–4 M strong acid with equimolar strong base. Before the equivalence point, the pH is determined by the untitrated strong acid, whereas after the equivalence point the concentration of excess strong base determines the pH. In an aqueous solution the concentration of H3O+ when the titration is 90% complete is
corresponding to a pH of 5.3. When the titration is 110% complete, the concentra- tion of OH– is
If the same titration is carried out in a nonaqueous solvent with a Ks of 1.0 x 10–20, the pH when the titration is 90% complete is still 5.3.
However, the pH when the titration is 110% complete is now
pH = pKs – pOH = 20.0 – 5.3 = 14.7
In this case the change in pH of
∆pH = 14.7 – 5.3 = 9.4
is significantly greater than that obtained when the titration is carried out in water. Figure 9.16 shows the titration curves in both the aque- ous and nonaqueous solvents. Nonaqueous solvents also may be used to increase the change in pH when titrating weak acids or bases (Figure 9.17).
Another parameter affecting the feasibility of a titration is the dis- sociation constant of the acid or base being titrated. Again, the solvent plays an important role. In the Brønsted–Lowry view of acid–base be- havior, the strength of an acid or base is a relative measure of the ease with which a proton is transferred from the acid to the solvent, or from the solvent to the base. For example, the strongest acid that can exist in water is H3O+. The acids HCl and HNO3 are considered strong because they are better proton donors than H3O+. Strong acids essentially donate all their protons to H2O, “leveling” their acid strength to that of H3O+. In a different solvent HCl and HNO3 may not behave as strong acids. When acetic acid, which is a weak acid, is placed in water, the dis- sociation reaction
CH3COOH(aq)+ H2O(l) < == == > H3O+(aq)+ CH3COO–(aq)
does not proceed to a significant extent because acetate is a stronger base than water and the hydronium ion is a stronger acid than acetic acid. If acetic acid is placed in a solvent that is a stronger base than water, such as ammonia, then the reaction
CH3COOH + NH3 < == == > NH4+ + CH3COO–
proceeds to a greater extent. In fact, HCl and CH3COOH are both strong acids in ammonia.
All other things being equal, the strength of a weak acid increases if it is placed in a solvent that is more basic than water, whereas the strength of a weak base in- creases if it is placed in a solvent that is more acidic than water. In some cases, how- ever, the opposite effect is observed. For example, the pKb for ammonia is 4.76 in water and 6.40 in the more acidic glacial acetic acid. In contradiction to our expec- tations, ammonia is a weaker base in the more acidic solvent. A full description of the solvent’s effect on a weak acid’s pKa or on the pKb of a weak base is beyond the scope of this text. You should be aware, however, that titrations that are not feasible in water may be feasible in a different solvent.