Primary standard substances, primary standard solution , secondary standard solution, end point, equivalence point, neutral point, indicators

Titration

Titration is an operation for the measurement of quantities of substances in solution by the method of 'volume analysis'. This process involves adding a solution of the reagent of known concentration (standard solution), taken in a burette (called 'titrant'), to an unknown solution (called analyte), taken in a flask known as titration flask. Titration is continued until the stoichiometric amounts of the two reagents (titrant and analyte) have been mixed. For example a titration of a strong acid such as HCl and a strong base such as NaOH involve the following reaction.

Equivalence point

The stoichiometric point corresponds to the neutralization of HCl and NaOH and is called equivalence point or end point. At this stage the number of moles of OH^- ions added to an acid solution becomes equal to number of moles of H⁺ ions present in the solution. Therefore, at equivalence point:

The simplest method for determining the equivalence point is to add a dye to the solution that shows change in colour at the equivalence point. Such dyes are called indicators. For example in the above titration of HCl with NaOH only NaCl is present in water at the equivalence point and the solution will have a pH of 7. Therefore, any indicator, which shows change in colour around pH 7 will be suitable for the titration.

Indicators

Indicators are substances, which indicate the completion (equivalence point or end point) of a chemical reaction by change in colour. For example, in volumetric analysis, during the titration of sodium hydroxide and hydrochloric acid (taken in the burette), phenolphthalein turns pink to colourless when the whole of sodium hydroxide has been neutralized by hydrochloric acid. All indicators show change in colour over some pH range, which varies considerably from one indicator to another.

Acid-Base Indicators

The amount of an acid (or a base), which is exactly equivalent chemically to the amount of some standard base (or an acid), is determined by an acid-base titration. The point of equivalence is called end point. The solution of a strong acid and strong base will be neutral at the end point and have a pH of 7 as they are strong electrolytes. However, if either the acid or the base is a weak electrolyte, the solution at the equivalence point will be either slightly alkaline (pH>7) or slightly acidic (pH<7).>⁺ ion concentration (i.e. pH), which depends upon the nature of the acid and the base and the concentration in the solution.

A large number of acid-base indicators are available which possess different colours according to H⁺ion concentration of the solution. These indicators do not change from predominantly 'acid' colour to predominantly 'base' suddenly and abruptly, but within a small interval of pH (generally about 2 pH units). This is called colour-change interval of the indicator. The position of the colour-change interval in the pH scale varies widely with different indicators. Therefore, for acid base titrations, an indicator is chosen which shows a distinct colour change at a pH close to that of equivalent point.

Colour changes of indicators with pH

Indicator	pH range	Colour Acid solution	Colour base solution
Cresol blue	1.2 — 1.8	Red	Yellow
Thymol blue	1.2 — 2.8	Red	Yellow
Methyl yellow	2.9 — 4.0	Red	Yellow
Methyl orange	3.1 — 4.4	Pink	Yellow
Methyl red	4.2 — 6.3	Red	Yellow
Litmus	5.0 — 8.0	Red	Blue
Bromothymol blue	6.0 — 7.6	Yellow	Blue
Phenol red	6.4 —8.2	Yellow	Red
Thymol blue (base)	8.1 — 9.6	Yellow	Blue
Phenolphthalein	8.3 — 10.0	Colourless	Blue
Thymolphthalein	8.3 — 10.5	Colourless	Blue
Alizarin yellow R	10.1 — 12.0	Blue	Yellow
Nitrarnine	10.0 — 13.0	Colourless	Orange/brown

Acid-base indicator behaviour

The mechanism of acid-base indicator was developed by Ostwald, which offered an explanation for the colour change with change in pH. All acid-base indicators are weak organic acids or bases, which possess different colours in unionized and ionized states. If the indicator acid is represented as HIn and has its ionized form In^-, then equilibrium in aqueous solution may be represented as:

For a basic indicator,

The unionized molecule has one colour while the ionized ion has another colour.

The equilibrium constant for a weak acid indicator may be written as:

where, K_Ind is known as indicator constant.

Writing the above equation in logarithmic form

Thus the colour of the indicator, which is determined by the ratio

Similarly, for a weak organic base indicator, In OH,

Taking logarithms

As the pH of the solution changes the colour of the indicator changes. This change is so minute that it goes unnoticed by the human eye. Therefore to see the colour change the concentration of one of the forms should predominate. The solution will have the acid colour (i.e. of HIn form) when the ratio [HIn] to [In^-] is approximately 10 and the alkaline colour (i.e. of In^-) when ratio [In^-] to [HIn] is approximately 10.

The acid colour will be visible when

Only basic colour will be visible when

Hence the colour change interval or indicator pH range is

Indicator pH range = -1 + pK_Ind to 1 + pK_Ind

Indicator pH range = pK_Ind ± 1

This means that the colour change interval of an indicator lies over approximately 2 pH units. Within this range, the change from one colour to the other another will be visible to human eye.

Mixed Indicators

With ordinary acid-base indicators the, colour change is not sharp and abrupt, but it extends over 2 units of pH. In order to have sharp colour change over a narrow and selected range of pH, a mixture of indicators are used. The pK_Ind values of these indicators are close together and the overlapping colours are complementary at an intermediate pH value.

For example, a mixture of 3 parts of phenolphthalein and 1 part of naphtholphthalein shows a change in colour from pale rose to violet at pH = 8.9. The mixed indicator is useful for the titration of phosphoric acid to diprotic stage.

Other examples:

Bromocresol green - methyl orange 4.3 (pH) Orange to blue green

Bromothymol blue - phenol red 7.5 (pH) Yellow to violet

Thymophthalein - phenolphthalein 9.9 (pH) Colourless to violet.