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 range. The colour change and the range of some common indicators are tabulated below


Table : Colour changes of indicators with pH

pH range
Acid solutionBase solution
Cresol red1.2 – 1.8RedYellow
Thymol blue1.2 – 2.8RedYellow
Methyl yellow2.9 – 4.0RedYellow
Methyl orange3.1 – 4.4PinkYellow
Methyl red4.2 – 6.3RedYellow
Litmus5.0 – 8.0RedBlue
Bromothymol blue6.0 – 7.6YellowBlue
Phenol red6.4 – 8.2YellowRed
Thymol blue (base)8.1 – 9.6YellowBlue
Phenolphthalein8.3 – 10.0ColourlessPink
Thymolphthalein8.3 – 10.5ColourlessBlue
Alizarin yellow R10.1 – 12.0BlueYellow
Nitramine10.8 – 13.0ColourlessOrange, Brown


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

Let be the equilibrium constant



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, , the colour visible will be yellow

, the colour visible will be red.

, the colour visible will be green.

In other words,

The colour visible will be red, when

The colour visible will be yellow, when

The colour visible will be green, when


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


 First change of colourMid point of changeColour change complete
[H+]10 KInKIn0.1 KIn
pHPKIn – 1PKInPKIn + 1


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


Tips & Tricks


?pH of boiling water is 6.5625. It does not mean that boiling water is not neutral. It is due to greater dissociation of H2O into H+ and OH.
?pH values of solutions do not give the exact idea of their relative strengths e.g., (i) A solution with pH = 1 had [H+] 100 times than that with pH = 3 and not 3 times. (ii) A 4 × 10–5 MHCl solution is twice concentrated as compared to 2 × 10–5 M HCl solution but pH values of these solutions are 4.4 and 4.7 respectively and not double.
?pH can be zero in 1M HCl or it can be negative for more concentrated solutions like 2M, 3M, 10M etc.
?At the temperature of the human body which is nearly 37°C, pH of neutral solution is 6.8.
?Buffer solutions have reserve acidity and reserve alkalinity.
?The greater the buffer capacity, the greater is its capacity to resist change in pH value.
?Buffers cannot withstand the addition of large amounts of acids or alkalies. The addition of 0.1mol per litre of [H+] or OH is about the maximum that any buffer can be expected to withstand.

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