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In acidimetry and alkalimetry, the choice of indicator depends on the hydrogen-ion concentration of the resulting salt. If a slightly alkaline salt like sodium acetate is formed at the equivalent point, an indicator is used which will show a change in colour on the alkaline side. Phenolphthalein, which changes colour for pH about 8 to 10, is therefore used in titrating acetic acid with sodium hydroxide. If a weak base like ammonia is being titrated with a strong acid, the resulting acid-reacting salt will require an indicator which changes colour on the acid side of neutrality. In this case, methyl-red, which changes at pH about 4 to 6, may be used. When a highly ionised neutral salt is formed, as in titrating a strong acid with a strong base, a slight excess of either solution makes a large change in pH, so that any indicator having a colour change between pH 3 and 10 will be satisfactory. Conversely, the titration of a very weak base with a very weak acid is seldom possible, on account of the relatively small change in pH. value at the end-point.
An important application of indicators is the determination of hydrogen ion concentration. An indicator is added to the unknown solution, and the tint compared with a series of reference solutions, containing the same indicator but varying H' concentrations. These standard solutions are "buffered" to definite pH values by mixtures such as borax and boric acid, or sodium acetate and acetic acid. The actual pH may be confirmed by hydrogen electrode measurements. The effect of salts, proteins, etc., on the indicator must be considered, and corrections applied.
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