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Indicators

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INDICATORS, in chemistry, are substances which are used, in virtue of colour changes which they undergo, in order to show when a certain reaction has reached completion, or to give in formation as to the acidity of a solution (see also CHEMISTRY: Analytical). Thus, if a solution of a neutral chloride is being titrated with silver nitrate, a few drops of potassium chromate solution are added; as soon as all the chloride has been precipi tated as silver chloride, the next drop of silver nitrate will form silver chromate (which is much less insoluble than silver chloride, but still fairly insoluble), and the precipitate will assume a faint reddish tinge whilst the liquid acquires a dull colour in place of the clear golden-yellow tint.

The indicators in general use may be divided into two classes (a) internal, and (b) external; for convenience, a large number of internal indicators used for determining acidity are dealt with as a separate class in HYDROGEN ION CONCENTRATION.

The potassium chromate in the above example is an internal indicator, because it is added to the solution undergoing titration with which it does not interfere. In titrating a solution of a ferrous salt by potassium dichromate, a very dilute solution of potassium ferricyanide is a useful indicator, but this has to be used externally, because, not only would it interfere chemically with the titration if used internally, but its colour changes could not be seen in the deep green solution. A drop of the titrated fluid is therefore withdrawn on a glass rod from time to time and mixed with a drop of the indicator solution on a white tile. As the titration approaches completion, the deep blue precipitate produced by the early drops gives place to a deep blue coloration, then pale blue, then green, then a slight dirty green, and finally no change at all, thus showing that the last trace of ferrous has been oxidized.

(a) Further examples of internal indicators are ferric alum in the titration of silver salts by thiocyanates (red coloration produced by the first drop in excess) ; starch solution in all iodo metric titrations (blue while the least trace of iodine is present) ; methylene-blue in titrations with Fehling's solution (decolorized as soon as all the latter solution is reduced) ; diphenylamine, with phosphoric acid, in dichromate–iron titrations (purplish-blue with first drop of dichromate in excess) ; potassium thiocyanate in iron–titanous chloride titrations (red whilst any ferric iron remains).

(b) Other external indicators are uranyl acetate in the titration of zinc salts by phosphates (brown precipitate as soon as the latter is in excess) ; potassium ferrocyanide in the titration of phosphates by uranyl acetate (brown precipitate with slightest excess) ; potassium iodide and starch in titrations involving ni trites or Fehling's solution (blue with a trace of either) ; ferrous thiocyanate with Fehling solution titrations (red until all the Fehling solution has been reduced) ; tannin solution in the titra tion of lead salts by ammonium molybdate (brownish-yellow colour with excess of the latter). (A. D. M.)

solution, potassium, titration, titrations and silver