Nonaqueous titration is the titration of substances dissolved in solvents other than water. It is the most common titrimetric procedure used in pharmacopoeial assays and serves a double purpose: it is suitable for the titration of very weak acids and very weak bases, and it provides a solvent in which organic compounds are soluble.
The most commonly used procedure is the titration of organic bases with perchloric acid in anhydrous acetic acid. These assays sometimes take some perfecting in terms of being able to judge the endpoint precisely.
The theory is that water behaves as both a weak acid and a weak base; thus, in an aqueous environment, it can compete effectively with very weak acids and bases with regard to proton donation and acceptance, as shown below:
The effect of this is that the inflection in the titration curves for very weak acids and very weak bases is small, because they approach the pH limits in water of 14 or 0 respectively, thus making endpoint detection relatively more difficult.
A general rule is that bases with pKa < 7 or acids with pKa > 7 cannot be determined accurately in aqueous solution.
Substances which are either too weakly basic or too weakly acidic to give sharp endpoints in aqueous solution can often be titrated in nonaqueous solvents. The reactions which occur during many nonaqueous titrations can be explained by means of the concepts of the Brønsted-Lowry theory. According to this theory an acid is a proton donor, i.e. a substance which tends to dissociate to yield a proton, and a base is proton acceptor, i.e. a substance which tends to combine with a proton. When an acid HB dissociates it yields a proton together with the conjugate base B of the acid:
Alternatively, the base B will combine with a proton to yield the conjugate acid HB of the base B, for every base has its conjugate acid and, every acid has its conjugate base.
It follows from these definitions that an acid may be either:
- an electrically neutral molecule, e.g. HCl, or
- a positively charged cation, e.g. C6H5NH3+, or
- a negatively charged anion, e.g. HSO4−.
A base may be either:
- an electrically neutral molecule, e.g. C6H5NH2, or
- an anion, e.g. Cl−.
Substances which are potentially acidic can function as acids only in the presence of a base to which they can donate a proton. Conversely basic properties do not become apparent unless an acid also is present.
Nonaqueous solvents used
Aprotic solvents are neutral, chemically inert substances such as benzene and chloroform. They have a low dielectric constant, do not react with either acids or bases and therefore do not favor ionization. The fact that picric acid gives a colorless solution in benzene which becomes yellow on adding aniline shows that picric acid is not dissociated in benzene solution and also that in the presence of the base aniline it functions as an acid, the development of yellow color being due to formation of the picrate ion.
Since dissociation is not an essential preliminary to neutralization, aprotic solvents are often added to 'ionizing' solvents to depress solvolysis (which is comparable to hydrolysis) of the neutralization product and so sharpen the endpoint.
Protophilic solvents are basic in character and react with acids to form solvated protons.
- HB + Sol. ⇌ Sol.H+ + B−
- Acid + Basic solvent ⇌ Solvated proton + Conjugate base of acid
A weakly basic solvent has less tendency than a strongly basic one to accept a proton. Similarly a weak acid has less tendency to donate protons than a strong acid. As a result, a strong acid such as perchloric acid exhibits more strongly acidic properties than a weak acid such as acetic acid when dissolved in a weakly basic solvent. On the other hand, all acids tend to become indistinguishable in strength when dissolved in strongly basic solvents owing to the greater affinity of strong bases for protons. This is called the leveling effect. Strong bases are leveling solvents for acids, weak bases are differentiating solvents for acids.
Amphiprotic solvents have both protophilic and protogenic properties. Examples are acetic acid and the alcohols. They are dissociated to a slight extent. The dissociation of acetic acid, which is frequently used as a solvent for titration of basic substances, is shown in the equation below:
- CH3COOH ⇌ H+ + CH3COO−
Here the acetic acid is functioning as an acid. If a very strong acid such as perchloric acid is dissolved in acetic acid, the latter can function as a base and combine with protons donated by the perchloric acid to form protonated acetic acid, an onium ion:
- HClO4 ⇌ H+ + ClO4−
- CH3COOH + H+ ⇌ CH3COOH2+ (onium ion)
Since the CH3COOH2+ ion readily donates its proton to a base, a solution of perchloric acid in glacial acetic acid functions as a strongly acidic solution.
When a weak base, such as pyridine, is dissolved in acetic acid, the acetic acid exerts its levelling effect and enhances the basic properties of the pyridine. It is possible, therefore, to titrate a solution of a weak base in acetic acid with perchloric acid in acetic acid, and obtain a sharp endpoint when attempts to carry out the titration in aqueous solution are unsuccessful.
- HClO4 + CH3COOH ⇌ CH3COOH2+ + ClO4−
- C5H5N + CH3COOH ⇌ C5H5NH+ + CH3COO−
- CH3COOH2+ + CH3COO− ⇌ 2CH3COOH
- Adding HClO4 + C5H5N ⇌ C5H5NH+ + ClO4−
- CH3COOH2+ + CH3COO− ⇌ 2CH3COOH
- C5H5N + CH3COOH ⇌ C5H5NH+ + CH3COO−
Titration of halogen acid salts of bases
The halide ions - chloride, bromide and iodide - are too weakly basic to react quantitatively with acetous perchloric acid. Addition of mercuric acetate (which is undissociated in acetic acid solution) to a halide salt replaces the halide ion by an equivalent quantity of acetate ion, which is a strong base in acetic acid.
- 2R.NH2.HCl ⇌ 2RNH3+ + 2Cl−
- (CH3COO)2Hg(undissociated) + 2Cl− → HgCl2 (undissociated) + 2CH3COO−
- 2CH3COOH2+ + 2CH3COO− ⇌ 4CH3COOH
The following indicators are in common use:
|Indicator||Color change||Color change||Color change|
|Crystal violet (0.5 per cent in glacial acetic acid)||violet||blue-green||yellowish-green|
|α-Naphtholbenzein (0.2 per cent in glacial acetic acid)||blue or blue-green||orange||dark-green|
|Oracet Blue B (0.5 per cent in glacial acetic acid)||blue||purple||pink|
|Quinaldine Red (0.1 per cent in methanol)||magenta||almost colorless|
The end point of most titrations is detected by the use of visual indicator but the method can be inaccurate in very dilute or colored solutions. However under the same conditions, a potentiometric method for the detection of the equivalence point can yield accurate results without difficulty. The electrical apparatus required consists of a potentiometer or pH meter with a suitable indicator and reference electrode. The other apparatus consists of a burette, beaker and stirrer.
The actual potential of the reference electrode need not be known accurately for most purposes and usually any electrode may be used provided its potential remains constant throughout the titration. The indicator electrode must be suitable for the particular type of titration (i.e. a glass electrode for acid-base reactions and a platinum electrode for redox titrations), and should reach equilibrium rapidly.
The electrodes are immersed in the solution to be titrated and the potential difference between the electrodes is measured. Measured volumes of titrant are added, with thorough (magnetic) stirring, and the corresponding values of emf (electromotive force) or pH recorded. Small increments in volume should be added near the equivalence point which is found graphically by noting the burette reading corresponding to the maximum change of emf or pH per unit change of volume. When the slope of the curve is more gradual it is not always easy to locate the equivalent point by this method. However, if small increments (0.1 cm³ or less) of titrant are added near the end point of the titration and a curve of change of emf or pH per unit volume against volume of titrant is plotted, a differential curve is obtained in which the equivalence point is indicated by a peak.
- Nonaqueous Acid-Base Titrations, Hardy Research Group, Department of Chemistry, The University of Akron