Ionic Equilibrium : Indicator theory

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

Indicator pH range Colour
Acid solution Base solution
Cresol red 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 Pink
Thymolphthalein 8.3 – 10.5 Colourless Blue
Alizarin yellow R 10.1 – 12.0 Blue Yellow
Nitramine 10.8 – 13.0 Colourless Orange, Brown

 

(1)     Theory of acid-base indicators : Two theories have been proposed to explain the change of colour of acid-base indicators with change in pH.

(i)     Ostwald’s Theory : According to this theory,

(a)     The colour change is due to ionisation of the acid-base indicator. The unionised form has different colour than the ionised form.

(b)     The ionisation of the indicator is largely affected in acids and bases as it is either a weak acid or a weak base. In case, the indicator is a weak acid, its ionisation is very much low in acids due to common H+ ions while it is fairly ionised in alkalies. Similarly if the indicator is a weak base, its ionisation is large in acids and low in alkalies due to common OH ions.

(ii)     Quinonoid theory : According to this theory,

(a)     The acid-base indicators exist in two tautomeric forms having different structures. Two forms are in equilibrium. One form is termed benzenoid form and other quinonoid form.

 

(b)     The two forms have different colours. The colour change is due to the interconversion of one tautomeric form into other.

(c)      One form mainly exists in acidic medium and other in alkaline medium.

Thus, during titration the medium changes from acidic to alkaline or vice-versa. The change in pH converts one tautomeric form into other and thus, the colour change occurs.

Phenolphthalein has benzenoid form in acidic medium and thus, it is colourless while it has quinonoid form in alkaline medium which has pink colour.

 

Methyl orange has quinonoid form in acidic solution and benzenoid form in alkaline solution. The colour of benzenoid form is yellow while that of quinonoid form is red.         

 

(2)     Selection of suitable indicator or choice of indicator : The neutralisation reactions are of the following four types,

(i)      A strong acid versus a strong base.

(ii)     A weak acid versus a strong base.

(iii)    A strong acid versus a weak base.       

(iv)    A weak acid versus a weak base.

In order to choose a suitable indicator, it is necessary to understand the pH changes in the above four types of 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.

In each case  of the acid (N/10) has been titrated against a standard solution of a base (N/10). Each curve becomes almost vertical for some distance (except curve of weak acid vs. weak base) and then bends away again. This region of abrupt change in pH indicates the equivalence point. For a particular titration, the indicator should be so selected that it changes its colour within vertical distance of the curve.

(i)      Strong acid Vs strong base : pH curve of strong acid (say HCl) and strong base (say NaOH) is vertical over almost the pH range 4 –10. So the indicators 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) are suitable for such a titration.

(ii)     Weak acid Vs strong base : pH curve of weak acid (say CH3COOH or oxalic acid) and strong base (say NaOH) is vertical over the approximate pH range 7 to 11. So phenolphthalein is the suitable indicator for such a titration.

(iii)    Strong acid Vs weak base : pH curve of strong acid (say HCl or H2SO4 or HNO3) with a weak base (say NH4OH) is vertical over the pH range of 4 to 7. So the indicators methyl red and methyl orange are suitable for such a titration.

(iv)    Weak acid vs. weak base : pH curve of weak acid and weak base indicates that there is no vertical part and hence, no suitable indicator can be used for such a titration.

(3)     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{\operatorname{Re}d}{\mathop{HIn(aq)}}\,\rightleftharpoons {{H}^{+}}(aq)+\underset{Green}{\mathop{I{{n}^{-}}(aq)}}\,

          Let KIn 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 = pKIn – 1

          The colour visible will be yellow, when pH = pKIn

          The colour visible will be green, when pH = pKIn + 1

Thus, our imaginary indicator will be red at any pH which just falls below pKIn – 1 and green at any pH which just exceeds pKIn + 1. The indicator changes its colour in the narrow pH range pKIn – 1 to pKIn + 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 pKIn of the indicator. The colour change of an indicator can, therefore, be summarised as,

 

  First change of colour Mid point of change Colour change complete
[H+] 10 KIn KIn 0.1 KIn
pH PKIn – 1 PKIn PKIn + 1

 

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