0
STANDARDIZATION OF ACID WITH POTASSIUM IODATE' PETER OESPER University of Pennsylvania School of Medicine, Philadelphia, Pennsylvania
TaE use of most of the common primary standards for acidimetry involves various disadvantages. Sodium carbonate may absorb water during weighing; borax must be recrystallized before it can be used; the salts of organic acids (sodium oxalate, for example) must be ignited to carbonate. Furthermore, it is difficult to make up and preserve standard solutions of carbonate or borax, because the alkaline solutions absorb carbon dioxide, or react with the glass of the containers. Potassium iodate has none of these drawbacks. Analytically pure preparations may be purchased and used without any further treatment except oven-drying. It is not hygroscopic, and its solutions are indefinitely stable. The standardization of acid with iodatetea is based on the reaction:
end point with dilute acid is no longer sharp, because the addition of a small amount of dilute acid does not change the pH of the solution enough to give a d e f i ~ t e change in the color of the indicator. These difficulties can be overcome by the use of a mixed indicator, which has a much sharper color change than a single indicator, and by the addition of a large excess of potassium iodide, which accelerates the reaction sufficiently, even a t pH 5, to make it possible to carry out the titration with only a few minutes' waiting a t the end point. Since thiosulfate also catalyzes the rea~tion,~.an excess of thiosulfate is a distinct advantage in the titration with acid more dilute than 0.05 N. The use of an excess of potassium iodide introduces one complication. Potassium iodide preparations usually contain several hundredths per cent of alkali, so that the addition of a large excess of potassium iodide would result in an overconsumption of acid. To avoid this error, the potassium iodide may be dissolved in water, and the solution neutralized beforehand with the acid to be standardized. The blank, for both indicator and reagents, amounts to about 0.15 ml. of 0.01 N acid. The following procedure is quite accurate for the standardization of strong acids whose concentration is 0.01 N or more. It is obviously not suitable for the standardization of weak acids, except in very concentrated solutions.
When a solution of iodate, to which potassium iodide, sodium thimlfate (in excess), and an acid-base indicator have been added, is titrated with acid, the reaction proceeds until the iodate is exhausted, and excess acid then causes the indicator to change color. The thiosulfate decolorizes the iodine formed, and makes it possible to observe the indicator end point. Near the end point, when the concentration of iodate has become small, the reaction proceeds very slowly, unless an excess of acid is present, since the rate of reaction is proportional to the concentration of iodate, and to the squares of the concentrations of hydrogen ion SOLUTIONS REQUIRED and of iodide iona Potassium Iodate. Standardized by direct weighing The usual suggestion, therefore, is to use an indicator and of approximately the same normality as the acid to with an end point distinctly on the acid side, such as be standardized (3.5 g. per liter is satisfactory for the methyl yellow. It is then possible to titrate directly standardization of 0.1 N acid). The use of solid potasto the end point (pH 3.5). Alternatively, an indicator sium iodate is not recommended, except for acid more like methyl red, which changes color a t pH 5, may be concentrated than 0.1 N, because of the low equivalent employed; it is then necessary to wait several minutes weight of the standard. for the reaction to go to completion, and to add more Indicator. Three parts 0.1 per cent bromcresol acid should the color revert to the alkaline shade as the green to 2 parts 0.1 per cent methyl red. This indicator last traces of iodate react. has a sharp transition from green to red a t pH 5. Either procedure is satisfactory when the acid is Potassium Iodide. Five-tenths of a gram per ml., 0.1 N or stronger, but difficulties arise in titrations neutralized to the mixed indicator with the acid to be with weaker solutions. If methyl red is used as indica- standardized. tor, the time required for the reaction of the last drop is excessive. and if methvl vellow is used. the indicator PROCEDURE error is t a d great, since farge excess o i dilute acid is Measure out enough potassium iodate solution to required to bring the pH down to 3.5. Moreover, the consume about 40 ml. of acid. Dilute to 50 ml. Add 2 ml. of potassium iodide solution, 9 drops ofindicator, The work reported here was carried out s t the University of and 800 mg. of sodium thiosulfate per 100 mg. of Maryland.
>
' EGGEET, J., Zeitschr. E l e k t r o b . , 30,501 (1924).
K O L ~ O F I. F ,M., S., J . Am. C h m . Sac., 48,1447 (1926). Wusnmx, S.,J. Phys. Chem., 8, 453 (1904).
C~RRIERE E., AND DALPLA, M., Cmnpt. rend., 205,1157 (1937).
588
NOVEMBER, 1949
potassium iodate. (In the standardization of acid more dilute than Q.05 N, use 1.2 g. of thiosulfate per 100 mg. of potassium iodate.) Titrate with the acid to be standardized, until the color changes from green through gray to pinkish, and remains on the pink side for three minutes (five minutes for acid more dilute than 0.05 N). A blank containing everything but potassium iodate should also be run. This procedure was tested experimentally on hydrochloric acid which had been standardized gravimetrically by precipitation of the chlorine as silver chloride,
589
and volumetrically with sodium carbonate and borax as primary standards. A total of eight determinations by these three methods gave a normality of 0.2000 += 0.0002. The results of four determinations with potassium iodate ranged from 0.2001 N to 0.2002 N. The acid was then accurately diluted t o normalities of 0.1000, 0.05000, and 0.01000. Four determinations of the first solution with potassium iodate ranged from 0.1000 to 0.1001; four determinations of the second, from 0.04996 to 0.05001; seven determinations of the third, from 0.00998 to 0.01000.
