Ionic Equilibrium


An indicator is a substance which is used to determine the end point in a titration. In acid-base titrations, organic substances (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:


pH range

Colour change

Methyl orange

Methyl red


Phenol red







Pink to yellow

Red to yellow

Red to blue

Yellow to red

Colourless to pink



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

1. 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.


Considering two important indicators phenolphthalein (a weak acid) and methyl orange (a weak base), Ostwald theory can be illustrated as follows:

Phenolphthalein: It can be represented as HPh. It ionises in solution to a small extent as:

HPh  ↔  H+ +  Ph-

Colourless        Pink


Applying law of mass action,

      K = [H+][Ph-]/[HpH]

The undissociated molecules of phenolphthalein are colourless while Ph-  ions are pink in colour. In presence of an acid the ionisation of HPh is practically negligible as the equilibrium shifts to left hand side due to high concentration of H+ ions. Thus, the solution would remain colourless. On addition of alkali, hydrogen ions are removed by OH- ions in the form of water molecules and the equilibrium shifts to right hand side. Thus, the concentration of Ph- ions increases in solution and they impart pink colour to the solution.

Let us derive Handerson equation for an indicator

HIn    + H2O ↔ H+3O  +   In-

'Acid form'               'Base form'


   Conjugate acid-base pair

     Kln = [ln][H+3O]/[HIn];              KIn = Ionization constant for indicator

     [H+3O] = KIn * [Hln]/ln-

     pH = -log10 [H+3O] =  -log10[Kln] - log10[Hln]/[ln-]

     pH = pKIn + log10[ln-]/[Hln]   (Handerson equation for indicator)

     At equivalence point

             [In-] = [HIn]  and  pH = pKIn


Methyl orange: It is a very weak base and can be represented as MeOH. It is ionized in solution to give Me+ and OH- ions.

        MeOH  ↔  Me+  + OH-

        Yellow     Red


Applying law of mass action,

                K = [Me+ ][OH- ]/[MeOH]

In presence of an acid, OH- ions are removed in the form of water molecules and the above equilibrium shifts to right hand side. Thus, sufficient Me+ ions are produced which impart red colour to the solution. On addition of alkali, the concentra­tion of OH" ions increases in the solution and the equilibrium shifts to left hand side, i.e., the ionisation of MeOH is practi­cally negligible. Thus, the solution acquires the colour of unionised methyl orange molecules, i.e., yellow.

This theory also explains the reason why phenolphthalein is not a suitableindicator for titrating a weak base against strong acid. The OH" ions furnished by a weak base are not sufficient to shift the equilibrium towards right hand side considerably, i.e., pH is not reached to 8.3. Thus, the solution does not attain pink colour. Similarly, it can be explained why methyl orange is not a suitable indicatorfor the titration of weak acid with strong base.

1.  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 the other quinonoid form.


(b) The two forms have different colors. The color change in due to the interconversation of one tautomeric form into other.

(c)  One form mainly exists in acidic medium and the 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 benziod 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 color of benzenoid form is yellow while that of quinoniod form is red.




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. (Fig. 10.1)

(ii) A weak acid versus a strong base. (Fig. 10.2)

(iii) A strong acid versus a weak base. (Fig. 10.3)

(iv) A weak acid versus a weak base. (Fig. 10.4)

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 theequivalence 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 titration curves of the above four types of neutralisation reactions are shown in Fig. 10.1, 10.2, 10.3 and 10.4.

In each case 25 mL of the acid (N/10) has been titrated against a standard solution of a base (N/10). Each titration curve becomes almost vertical for somedistance (except curve 10.4) 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 HCI) 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 3.2-4.5) are suitable for such a titration.

(ii) Weak acid vs. weak base:

pH curve of weak acid (say CH3COOH of 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.



(iii) 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.

Titration of soluble carbonate with strong acid.

pH curve of sodium carbonate with HCI shows two inflec­tion points (Fig. 10.5). First inflection point (pH 8.5) indicates conversion of carbonate into bicarbonate.

Na2CO3 + HCI  → NaHCO3 + NaCl

As the inflection point lies in the pH range 8 to 10, phenolphthalein can be used to indicate the above conversion. The second inflection point (pH 4.3) indicates the following reaction:

NaHCO3 + HCI  → NaCl + CO2 + H2O

As the point lies between 3 to 5, methyl orange can be used.

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