Submitted by ChemPRIME Staff on Thu, 12/16/2010 - 14:55

We mentioned that in most titrations it is necessary to add an indicator which produces a sudden color change at the equivalence pointThe point in a titration at which the amount of one reactant being added stoichiometrically matches the amount of another reactant initially present. The end point should match the equivalence point as closely as possible..
Colorless phenolphthalein (3,3-bis(4-hydroxyphenyl)-2-benzofuran-1-one) at low-pH
A typical indicator for acidIn Arrhenius theory, a substance that produces hydrogen ions (hydronium ions) in aqueous solution. In Bronsted-Lowry theory, a hydrogen-ion (proton) donor. In Lewis theory, a species that accepts a pair of electrons to form a covalent bond.-baseIn Arrhenius theory, a substance that increases the concentration of hydroxide ions in an aqueous solution. In Bronsted-Lowry theory, a hydrogen-ion (proton) acceptor. In Lewis theory, a species that donates a pair of electrons to form a covalent bond. titrations is phenolphthalein, HC20H13O4. Phenolphthalein, whose structure is shown below, is a colorless weak acidAn acid that ionizes only partially in a given solvent. (Ka = 3 × 10–10 mol dm–3). Its conjugate baseThe base formed when an acid releases a hydrogen ion (proton)., C20H13O4 has a strong pinkish-red color. In order to simplify, we will write the phenolphthalein moleculeA set of atoms joined by covalent bonds and having no net charge. as HIn (protonated indicator) and its pink conjugate base as In. In aqueous solution, phenolphthalein will present the following equilibriumA state in which no net change is occurring, that is, in which the concentrations of reactants and products remain constant; chemical equilibrium is characterized by forward and reverse reactions occurring at the same rate.

HIn + H2O \rightleftharpoons In + H3O+      (1)

According to Le Chatelier’s principle, the equilibrium expressed in equation (1) will be shifted to the left if H3O+ is added. Thus in a strongly acidic solution we expect nearly all the pink In to be consumed, and only colorless HIn will remain. On the other hand, if the solution is made strongly basic, the equilibrium will shift to the right because OH ions will react with HIn molecules, converting them to In. Thus the phenolphthalein solution will become pink.

Clearly there must be some intermediateIn chemical kinetics, a species that is formed in an early step in a reaction mechanism and then consumed in a later step; evidence of existence of an intermediate may be important for the interpretation of a rate law. situation where half the phenolphthalein is in the acid form and half in the colored conjugate-base form. That is, at some pH

      [HIn] = [In]

This intermediate pH can be calculated by applying the Henderson-Hasselbalch equation to the indicator equilibrium:

\text{pH}=\text{p}K_{a}\text{ + log}\frac{[\text{ In}^{-}]}{[\text{ HIn }]}

Thus at the point where half the indicator is conjugate acidThe acid formed when a base accepts a hydrogen ion (proton). and half conjugate base,

\text{pH}=\text{p}K_{a}\text{ + log 1} = \text{p}K_{a}\,

For phenolphthalein, we have

\text{pH}=\text{p}K_{a}=-\text{log(3 }\times \text{ 10}^{-10}\text{)}=\text{9.5}

so we expect phenolphthalein to change color in the vicinity of pH = 9.5.

Fig. 1 Color change of phenolphthalein with respect to pH

Fig. 2 Phenolphtalein at different pH, notice the distinctive color at pH 8-12

The way in which both the color of phenolphthalein and the fraction present as the conjugate base varies with the pH is shown in detail in Fig. 1. The change of color occurs over quite a limited range of pH―roughly pKa ± 1. In other words the color of phenolphthalein changes perceptibly between about pH 8.3 and 10.5. Observe the actual color change for this indicator in Fig. 2. Other indicators behave in essentially the same way, but for many of them both the acid and the conjugate base are colored. Their pKa’s also differ from phenolphthalein, as shown in the following table. The indicators listed have been selected so that their pKa values are approximately two units apart. Consequently, they offer a series of color changes spanning the whole pH range.

Properties of Selected Indicators

   Name pKa Effective pH range Acid form Basic form
  Thymol blue     1.6   1.2 - 2.8     Red     Yellow
  Methyl orange     4.2   3.1 - 4.4     Red     Orange
  Methyl red     5.0   4.2 - 6.2     Red     Yellow
  Bromothymol blue     7.1   6.0 - 7.8     Yellow     Blue
  Phenophthalein     9.5   8.3 - 10.0     Colorless     Red
  Alizarin yellow   11.0 10.1 - 12.4     Yellow     Red

Indicators are often used to make measurements of pH which are precise to about 0.2 or 0.3 units. Suppose, for example, we add two drops of bromothymol blue to a sample of tap water and obtain a green-blue solution. Since bromothymol blue is green at a pH of 6 and blue at a pH of 8, we conclude that the pH is between these two limits. A more precise result could be obtained by comparing the color in the tap water with that obtained when two drops of indicator solution are added to buffer solutions of pH 6.5 and 7.5.

EXAMPLE 1 What indicator, from those listed in the table, would you use to determine the approximate pH of the following solutions:

a) 0.1 M CH3COONa (sodium acetate)

b) 0.1 M CH3COOH (acetic acid)

c) A bufferA solution to which strong acid or strong base can be added without significant change in pH; a solution containing a weak acid and its conjugate base or a weak base and its conjugate acid. mixtureA combination of two or more substances in which the substances retain their chemical identity. of sodium acetate and acetic acid

d) 0.1 M NH4Cl (ammonium chloride)


a) A solution of sodium acetate will be mildly basic with a pH of 9 or 10. Phenolphthalein would probably be best.

b) A solution of acetic acid, unless very dilute, has a pH in the vicinity of 3. Both thymol blue and methyl orange should be tried.

c) Since Ka for acetic acid is 1.8 × 10–5 mol dm–3, we can expect this buffer to have a pH not far from 5. Methyl red would be a good indicator to try.

d) Since NH4+is a very weak acid, this solution will be only faintly acidic with a pH of 5 or 6. Again methyl red would be a good indicator to try, though bromothymol blue is also a possibility.

If a careful choice of both colors and pKa is made, it is possible to mix several indicators and obtain a universal indicator which changes color continuously over a very wide pH range. With such a mixture it is possible to find the approximate pH of any solution within this range. So-called pH paper is impregnated with one or several indicators. When a strip of this paper is immersed in a solution, its pH can be judged from the resulting color.