Theory of Indicators

Generally the theory based on that all indicators are either weak acids or weak bases in which the color of the ionized form is different from the color before dissociation.

They were first systematically employed in analytical chemistry by Robert Boyle, who used the aqueous extracts of the coloured principles present in red-cabbage, violets and cornflowers. The indicator most in use to-day is litmus, whose solution is turned red by an acid, and blue by an alkali. Several synthetic indicators are employed in acidimetry and alkalimetry. The choice is not altogether arbitrary, for experiments have shown that some are more suitable for acidimetry, while others are only applicable in alkalimetry; moreover, the strength of the acids and bases employed may exert a considerable influence on the behavior of the indicator.

litmus paper

Theory of Indicators 
 The ionic theory of solutions permitted the formulation of a logical conception of the action of indicators by W. Ostwald which for many years held its ground practically unchallenged; and even now the arguments originally advanced hold good, except for certain qualifications rendered necessary by more recent research. In the language of the ionic theory, an acid solution is one containing free hydrions, and an alkaline solution is one containing free hydroxidions. A neutral solution contains hydrions and hydroxidions in equal concentration; this is a consequence of the fact that pure water itself undergoes a certain dissociation, and several different methods show that in the purest water obtainable the concentration of the free hydrions and hydroxidions is 107 at 24°. Moreover, the law of mass-action (see Chemical Action) demands that the product of the concentrations of the hydrions and hydroxidions in any solution is constant at a given temperature, and we see from the above values that this constant is 1014. It follows, therefore, that the acidity or alkalinity of any solution can be expressed both in terms of hydrion or hydroxidion concentration. Many researches have been directed to classify acid and alkaline solutions according to the concentration of the hydrion. Conductivity determinations show that the maximum concentration of hydrion occurs in 5.8 - N nitric acid, where it has a value of about 2 -N, and the minimum occurs in 6.7N potassium hydroxide, where its value is 5 X I 015, that of the hydroxidion being about 2 - N. These figures apply to a temperature of 24 °. Bearing in mind the concentration of the ions in a neutral solution, it is seen that a scheme of seven grades of "neutrality," differing by successive powers of ten, may be formulated. The concentration of hydrion and hydrox idion in any solution may be determined by several independent methods, and it is therefore a simple matter to prepare solutions of definite ionic concentrations and to test these with the object of obtaining a list of indicators according to their sensitiveness. It is found that litmus responds to concentrations of io-- 6 H and 106 0H', a result which shows this dye to be the best indicator of true neutrality. Methyl orange responds to between - 4H and I 05 H - para-nitrophenol to between io - 5 H- and 106 Hand phenolphthalein to between r05 0H' and t06 0W. Salm (Zeit. Elektrochem., 1904, 10, p. 341) gives a list of twenty-seven indicators classified on this principle. Other papers bearing on this subject are Friedenthal, ibid., p. 113; Salessky, ibid., p. 204; Fels, ibid., p. 208; Scholtz, ibid., p. 549; M. Handa, Ber., 1909, 4 2, p. 3179.

The actual mechanism by which the indicator changes colour with varying concentrations of hydrion or hydroxidion is now to be considered. Ostwald formulated his ionization theory which assumes the change to be due to the transition of the non-dissociated indicator to the ionized condition, which are necessarily of different colours. On this theory, an indicator must be weakly basic or acid, for if it were a strong acid or base high dissociation would occur when it was in the free state, and there would be no change of colour when the solution was neutralized. Take the case of a weakly acid indicator such as phenolphthalein. The presence of an acid depresses the very slight dissociation of the indicator, and the colour of the solution is that of the non-dissociated molecule. The addition of an alkali, if it be strong, brings about the formation of a salt of phenolphthalein, which is readily ionized, and so reveals the intense red coloration of the anion; a weak base, however, fails to give free ions. An acid indicator of medium strength is methyl orange. When free this substance is ionized and the solution shows an orange colour, due to a mixing of the red of the non-dissociated molecule and the yellow of the ionized molecule. Addition of hydrions lessens the dissociation and the solution assumes the red colour, while a base increases the dissociation and so brings about the yellow colour. If the alkaline solution be titrated with a strong acid, the hydrions present in a very small amount of the acid suffices to reverse the colour; a weak acid, however, must be added in considerable excess of the quantity properly required to neutralize the solution, owing to its weak dissociation. This indicator is therefore only useful when strong acids are being dealt with, while it's strongly acid nature renders it serviceable for both strong and weak bases.


pH indicator 
 A pH indicator is a halochromic chemical compound that is added in small amounts to a solution so that the pH (acidity or alkali nity) of the solution can be determined easily. Hence a pH indicator is a chemical detector for hydronium ions (H3O+) (or Hydrogen ions (H+) in the Arrhenius model). Normally, the indicator causes the color of the solution to change depending on the pH. Solutions with a pH value above 7.0 are alkali, and solutions with a pH value below 7.0 are acidic. Solutions with a pH value of 7.0 are neutral 

Theory 
 pH indicators themselves are frequently weak acids or bases. When introduced into a solution, they may bind to H+ (Hydrogen ion) or OH- (hydroxide) ions. The different electron configurations of the bound indicator causes the indicator's color to change, which allows the pH to be determined by the different colors. 

Application 
pH indicators are frequently employed in titrations in analytic chemistry and biology experiments to determine the extent of a chemical reaction. Because of the subjective determination of color, pH indicators are susceptible to imprecise readings. For applications requiring precise measurement of pH, a pH meter is frequently used.

There are several common laboratory pH indicators. Indicators usually exhibit intermediate colors at pH values inside the listed transition range. For example, phenol red exhibits an orange color between pH 6.8 and pH 8.4. The transition range may shift slightly depending on the concentration of the indicator in the solution and on the temperature at which it is used.

Naturally occurring pH indicators 
 Many plants or plant parts contain chemicals from the naturally-colored anthocyanin family of compounds. They are red in acidic solutions and blue in basic. Extracting anthocyanins from red cabbage leaves or the skin of a lemon to form a crude acid-base indicator is a popular introductory chemistry demonstration.

Anthocyanins can be extracted from a multitude of colored plants or plant parts, including from leaves (red cabbage); flowers (geranium, poppy, or rose petals); berries (blueberries, blackcurrant); and stems (rhubarb). An exhaustive list would be beyond the scope of this article

Sources :

www.thefreedictionary.com

www.1911encyclopedia.org/Indicator

en.wikipedia.org/wiki/PH_indicator