JEE Main & Advanced Chemistry Equilibrium / साम्यावस्था Indicators

An indicator is a substance, which is used to determine the end point in a titration. In acid-base titrations, organic substance (weak acids or weak bases) are generally used as indicators. They change their colour within a certain \[pH\] range. The colour change and the \[pH\] range of some common indicators are tabulated below

Colour changes of indicators with pH

Two theories have been proposed to explain the change of colour of acid-base indicators with change in pH.

(i) Ostwald’s Theory  (ii) Quinonoid theory

(1) Selection of suitable indicator or choice of indicator : In order to choose a suitable indicator, it is necessary to understand the pH changes in the titrations. The change in pH in the vicinity of the equivalence point is most important for this purpose. The curve obtained by plotting pH as ordinate against the volume of alkali added as abscissa is known as neutralisation or titration curve. The suitable indicators for the following titrations are,         

(i) Strong acid Vs strong base : Phenolphthalein (pH range 8.3 to 10.5), methyl red (pH range 4.4 – 6.5) and methyl orange (pH range 3.2 to 4.5).

(ii) Weak acid Vs strong base : Phenolphthalein.

(iii) Strong acid Vs weak base : Methyl red and methyl orange.

(iv) Weak acid vs. weak base : No suitable indicator can be used for such a titration.

Reason for use of different indicators for different systems : Indicators are either weak acids or weak bases and when dissolved in water their dissociated form acquires a colour different from that of the undissociated form. Consider a weak acid indicator of the general formula HIn, where in represents indicator. The equilibrium established in aqueous solution will be

\[\underset{\text{Red}}{\mathop{HIn(aq.)}}\,\] ? \[{{H}^{+}}(aq.)+\underset{\text{Green}}{\mathop{I{{n}^{-}}(aq.)}}\,\]

Let \[{{K}_{In}}\] be the equilibrium constant

\[{{K}_{In}}=\frac{[{{H}^{+}}][I{{n}^{-}}]}{[HIn]}\] or \[\frac{[HIn]}{[I{{n}^{-}}]}=\frac{[{{H}^{+}}]}{{{K}_{In}}}\]

The human eye can detect the change in colour if the ratio of the two forms of indicator ranges between 0.1 to 10.

If, \[\frac{[HIn]}{[I{{n}^{-}}]}=1.0\], the colour visible will be yellow

\[\frac{[HIn]}{[I{{n}^{-}}]}=10\], the colour visible will be red.

\[\frac{[HIn]}{[I{{n}^{-}}]}=0.1\], the colour visible will be green.

In other words,

The colour visible will be red, when \[pH=p{{K}_{In}}-1\]

The colour visible will be yellow, when \[pH=p{{K}_{In}}\]

The colour visible will be green, when \[pH=p{{K}_{In}}+1\]

Thus, our imaginary indicator will be red at any \[pH\] which just falls below \[p{{K}_{In}}-1\] and green at any \[pH\] which just exceeds \[p{{K}_{In}}+1\]. The indicator changes its colour in the narrow \[pH\] range \[p{{K}_{In}}-1\] to \[p{{K}_{In}}+1\] from red to (red-yellow, yellow, yellow-green) green. We can therefore use this indicator to locate this narrow \[pH\] range. In other words, in order to use the indicator effectively in this range, we should have a solution for which \[pH\] is very near to \[p{{K}_{In}}\] of the indicator. The colour change of an indicator can, therefore, be summarised as,

It is for this reason that we use different indicators for different systems.