Acid-base equilibria
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What is the science behind methyl orange indicator? What colour changes occur during titration with a strong acid and a weak alkali?
Jenkin replies ...
All acid-base indicators are weak acids, so just a small proportion of their molecules ionise in aqueous solution. We can show this as an equation:
HIn <=> H+ + In–
HIn represents the indicator which is a weak acid, and In– the ion it forms when it ionises. HIn and In– have different colours: to take the example of methyl orange, HIn is red and In– is yellow.
Now, we said it was a weak acid, so it will be mostly in the HIn form and therefore coloured red. But if we add alkali (such as NaOH) to methyl orange, this will react with the H+ ions and then more HIn will ionise so as to replace them. So more In– will be formed and eventually the indicator will appear yellow. This will show up when the In– concentration is about 10 times greater than the HIn concentration.
The reaction is reversible, so you can start with the indicator having either colour.
All indicators change colour over a particular pH range and this has to be considered when we are choosing an indicator for a particular titration. When a strong acid and a strong alkali are involved, almost any of the common indicators can be used.
Strong acid/weak alkali titrations: at the endpoint of such a titration, the pH of the solution is likely to be in the region of pH 3 to 5, which corresponds to the pH range over which methyl orange changes colour, so it is OK to use it. But phenolphthalein changes over the range pH 8 to 10, and a strong acid/weak alkali titration will never have a pH in this region.
You will find some discussion of this subject in the Nuffield Chemistry Students' book, pages 342 -343 in Topic 14.
You will find a fuller list of indicators and their pH ranges in the Nuffield Book of Data, Table 6.6 on page 123. You will also find that most other A-level (but probably not AS) textbooks deal with this subject.
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updated: 06 November 2003
