SC VI B. Titrimetric Analysis
This Supplementary Chapter provides general guidance regarding the practice of titration in quantitative pharmaceutical analysis. It is not intended as an exhaustive guide to the underlying theory.
Basic concepts
Titration is the process of dissolving an analyte and reacting with another species in a solution of known concentration (titrant). Once the reaction between the analyte and the titrant is well understood, the stoichiometry is characterised and the point at which the titrant amount is equivalent to the analyte amount (the equivalence point) is known, the exact quantity of the analyte can be determined by simple calculation.
Acid-base titrations The determination is based on the neutralisation reactions that occurs between an acid and a base. A measured volume of an acid or base of known concentration is reacted with a sample to the equivalence point. The relative acidity (basicity) of an aqueous solution can be determined using the relative acid (base) equivalents. An acid equivalent is equal to one mole of H+ ions. Similarly, a base equivalent is equal to one mole of OH– ions. Some acids and bases are polyprotic, which means that each mole of the acid or base is capable of releasing more than one acid or base equivalent. Examples of monoprotic acids are hydrochloric acid (HCl) and nitric acid (HNO3). Examples of polyprotic acids are sulfuric acid (H2SO4) and phosphoric acid (H3PO4).
When the solution of known concentration and the solution of unknown concentration are reacted to the point where the number of acid equivalents equals the number of base equivalents (or vice versa), the equivalence point is reached. The equivalence point of a strong acid or a strong base will occur at pH 7. For weak acids and bases, the equivalence point need not occur at pH 7. There will be several equivalence points for polyprotic acids and bases.
Redox titrations The determination is based on the transfer of electrons between a donor (reducing agent) and an acceptor (oxidising agent). During the reaction, the oxidising ion, whether it is the analyte or the titrant, is reduced by gaining one or more electrons and the reducing ion is oxidised, losing one or more electrons.
These reactions are less common than acid/base reactions but involve a wider range of titrants including oxidising agents (such as iodine, potassium dichromate, potassium permanganate solutions, cerium(iv) salts, hydrogen peroxide, oxidised chlorine (for example ClO–, ClO), and reducing agents (such as sodium thiosulfate, oxalic acid, ammonium iron(ii) sulfate (“Mohr’s salt”), hydrogen peroxide, phenylarsine oxide).
Complexometric titrations The determination is based on the formation of a complex usually between a chelating agent and a metal cation. The most frequent use of complex reactions is to determine the concentration of divalent cations such as calcium, magnesium, copper, lead, zinc and cadmium as well as other cations such as aluminium.
The most commonly used complexing reagents are ethylenediaminetetraacetic acid (EDTA) and ethylenebis(oxyethylenenitrilo)tetracetic acid (EGTA).
Although these reactions are easy to perform, it is necessary to work within a well-defined pH interval.
Precipitation titrations The determination is based on the formation of an insoluble salt under certain conditions. During a titration, the end of the precipitation reaction means that excess titrant is present and a coloured complex appears immediately. They are performed in slightly acidic conditions (pH about 4.5) and solvents such as ethanol or acetone may be added to reduce the solubility of the precipitate.
The most frequent use of precipitation reactions in analytical chemistry is the titration of halides (in particular chloride ions) with silver ions.
Direct titrations
In a direct or forward titration the analyte reacts directly with the titrant. For example, a known volume of a solution of unknown acidity may be titrated with a base of known concentration until complete neutralization has occurred. This point is called the equivalence point and is generally determined by observing a colour change in an added indicator such as phenolphthalein. From the volume and concentration of added base and the volume of acid solution, the unknown concentration of the solution before titration can be determined.
Titrations can also be used to determine the number of acidic or basic groups in an unknown compound. A specific weight of the compound is titrated with a known concentration of acid or base until the equivalence point has been reached. From the volume and concentration of added acid or base and the initial weight of the compound, the equivalent weight, and thus the number of acidic or basic groups, can be determined.
Examples of direct titrations include the Assay for Fusidic Acid and the Assay for Ketoconazole.
Back titrations
The term back titration is used when a titration is done “backwards”; instead of titrating the original analyte, a volumetric solution of a reagent is added to the solution to react with the analyte, then the excess reagent is titrated.
Back titrations are useful if the end point of the reverse titration is easier to identify than the end point of the normal titration. They are also useful if the reaction between the analyte and the titrant is very slow. They may also be used where direct titrations are unsuitable for technical reasons, including: when the sample is not soluble in water, or when the sample contains impurities that interfere with forward titration.
Examples of back titrations include the Assay for Aspirin Tablets and the Assay for Cloxacillin Benzathine.
Measuring the equivalence point
The equivalence point may be determined in a number of ways. The most common is the use of an indicator that shows a colour change at the equivalence point.
Alternatively, the pH or potential difference of the solution can be plotted against the amount of added acid or base on a graph; such a plot is called a titration curve and is usually sigmoid (S-shaped), with the inflection point where the curve changes direction corresponding to the equivalence point. From the pH at the equivalence point, the dissociation constant of the acidic or basic group can be determined. If a compound contains several different acidic or basic groups, the titration curve will show several sigmoid-shaped curves and the dissociation constant of each group can be obtained from the pH at its corresponding equivalence point.
Indicators
Indicators are used to provide a visual determination of the end point of a reaction. This may be by a change in colour or the formation of a precipitate.
Because they are used in low concentrations, indicators do not appreciably alter the equivalence point. Sometimes the volume difference (error) is ignored; in other cases a correction factor may be applied. Often a sensitivity test may be used as a suitability check before an indicator is used.
Examples of indicators where sensitivity tests are specified include Methyl Orange Solution, Phenolphthalein Solution, Starch Solution and Thymol Blue Solution.
Acid/base Indicators Acid-base indicators are weak acids or weak bases. The undissociated form of the indicator has a different colour to the dissociated form. Indicators do not change colour from pure acid to pure alkali at a specific hydrogen ion concentration; rather the colour change occurs over a range of hydrogen ion concentrations. This range is often termed the colour change interval and is expressed as a pH range.
Acid-base indicators include Bromocresol Green, Cresol Red, Crystal Violet, Methyl Orange, Phenolphthalein and Thymol Blue.
An example of this type of titration is the Assay for Aspirin Tablets.
Redox Indicators Redox indicators undergo a definite colour change at a specific electrode potential. As they have different colours in their oxidised and reduced states, their colour changes according to the redox potential of the solution. As for those used in pH measurement, these indicators have specific colour change intervals expressed in mV. They may be dependant or independent of pH.
Redox indicators include Nitrophenanthroline, Diphenylamine, Indigo Carmine, Methylene Blue and Ferroin.
An example of this type of titration is the Degree of unsaturation test in the monograph for Undecenoic Acid.
Complexometric indicators Complexometric indicators are ionochromic dyes that undergo a definite colour change in the presence of specific metal ions. They form a weak complex with the ions present in the solution, which has significantly different colour than the uncomplexed form. They are also called metallochromic indicators.
Complexometric indicators are water-soluble organic molecules, including Eriochrome Black T, Xylenol Orange, Murexide, Eriochrome Blue SE, Methylthymol Blue and Naphthol Green B.
An example of this type of titration is the Assay in the monograph for Hydrotalcite Tablets.
Precipitation indicators Precipitation indicators precipitate from solutions in a readily visible form at or near the equivalence point of a titration.
Precipitation indicators include Fluorescein and Eosin.
The colour change intervals for several common indicators are given in Supplementary Chapter VI C.