CHLOROMETRY A
Titrimetric Procedure A v a i l a b l e for Microanalysis NATHAN I. GOLDSTONE AND MORRIS B. JACOBS Department of Health, City of N e w York, N. Y. PREPARATION OF STANDARD SODIUM HYPOCHLORITE SOLUTION
Standard sodium hypochlorite solution is recommended as a general titrimetric reagent for microanalytical work. The solution must b e made with sufficient excess of sodium hydroxide so that the p H i s about 19.5 and must b e stored in brown, glass-stoppered bottles. Under these conditions, sodium hypochlorite solution is remarkably stable and compares favorably with other titrimetric reagents. With the use OF this reagent, chlorometric determinations can b e made at room temperature with ease, precision, accuracy, and a sharpness of the end point that leaves nothing to be desired in an analytical reagent.
Undoubtedly one reason for the lack of enthusiasm among analysts for the use of sodium hypochlorite as a titrimetric reagent is the apparent difficulty of preparing such solutions. The authors discarded the cumbersome method of Jellinek and Kresteff ( 5 ) and prepared a standard solution in the following simple manner. Transfer 8.0 ml. of a commercial preparation of sodium hypochlorite solution containing 5% of available chlorine to a glassstoppered brown-glass bottle, and dilute with water to about 2 liters. If necessary, add sufficient sodium hydroxide (1 gram) to raise the pH to about 12.5, the optimum pH for stability. To ascertain if the proper pH has been reached, the customary colorimetric methods for the determination of pH in the range 12 to 14 may be used. The authors used Chlorox as the source of sodium hypochlorite. Obtain the titer of the solution by titration against a primary standard of sodium arsenite made as follows: Weigh 0.2473 gram of arsenious oxide (arsenic trioxide, Aslo,, Kational Bureau of Standards) and dissolve in 25 ml. of 10% sodium hydroxide solution. Transfer to a 1-liter volumetric flask, make slightly acid with sulfuric acid (1 to 6), and dilute with water to 1 liter. This solution is 0.005 N .
IT
I S a popular misconception that because sodium hypochlorite is a highly reactive substance it is too unstable to be used as a standard titrimetric reagent. Because of this belief, it has almost never been used for this purpose. “Chlorometry” is the term used to designate the quantitative estimation of various substances by use of standard hypochlorite solution, in a manner entirely analogous to iodometry and bromometry. Although chlorometric determinations have not achieved any degree of practical application, there are references to the use of hypochlorite as a macroanalytical reagent in the literature.
The solution of sodium hypochlorite made as directed above is generally somewhat stronger than 0.005 N . Its exact titer can be determined by titration against the standard arsenite solution. Its normality may be adjusted to exactly 0.005 N by the usual procedure.
As early as 1824, Gay-Lussac estimated the strength of chlorinated lime by rmitting a fine sus nsion of chlorinated lime to flow from a G e t into a hydroc%oric acid solution of an arsenite and obtained the end point of the reaction by using indigo as the indicator. Deniges (9) using a similar modification turned to potassium bromide as the indicator. However, Lunge-Berl (9) showed both of these determinations to be rather inaccurate. In some respects, chlorometry has as venerable a history as iodometry, introduced by Bunsen and Schwarz in 1853, the bromate methods of Koppeschaar (8) and the bromometric methods of Manchot (IO). Jackson and Parsons ( 4 ) advocated the use of sodium chlorite as a volumetric oxidizing a ent. Jellinek and Kresteff (6),continuing a series o f studies of newer methods in volumetric analysis, presented chlorometry as a valuable volumetric aid. They pre ared a standard solution of sodium hypochlorite by passing cborine gas from a small tank of chlorine into N sodium hydroxide solution until the solution was approximately 0.13 N with respect to sodium hypochlorite content. They kept this solution in a clear transparent bottle connected to a buret, using rubber connections, and found that with an excess of alkali the titer of the standard sodium hypochlorite solution kept surprisingly constant. Thus, their ori inal titer, using the potassium iodide-hydrochloric acid mettod of estimation, was 0.1350 to 0.1351 N . After 7 days, the strength of this solution was practically unaltered, for their estimations gave the normality a value of 0.1351 to 0.1352. After an additional 10 days the titer had been reduced to 0.1330 to 0.1328. They atdbuted this loss of about 1.5% to the fact that they had made no attempt t o shield the solution from sunlight. Jellinek and Kuhn (6) prepared a standard solution of sodium hypochlorite in the manner described by Jellinek and Kresteff and found it practically constant for a number of weeks. Kolthoff and Stenger ( 7 ) found that “H.T.H.” calcium hypochlorite yielded stable solutions. They added excess bromide to the sample to be titrated, so that in their titrations the added hypochlorite actually behaved as hypobromite. They standardized their hypochlorite solutions by titration against arsenic trioxide in acid or weakly alkaline solution, using Bordeaux as
TITRATION PROCEDURE
To estimate the strength of the sodium hypochlorite solution the following simple procedure may be used. Transfer a known aliquot of standard arsenite solution to a
125-ml. Erlenmeyer flask or a 150-ml. beaker: a 4-ml. aliquot
if a microburet is to be used for the standard hypochlorite solution and a 5-ml. aliquot if a semimicroburet is to be used. A standard solution of tartar emetic, [potassium antimonyl tartrate, K(SbO)C4H406.SHzO,] containing 1 mg. of antimony per 10 ml. of solution may also be used. Add 5 ml. of concentrated hydrochloric acid and adjust the volume of the solution to 35 to 40 ml. by adding distilled water. Fill a microburet or semimicro buret with the standard hypochlorite solution. Add 1 drop of 0.05’% methyl orange indicator solution to the test solution and titrate directly with the sodium hypochlorite solution. Add another drop of methyl orange indicator solution near the end point and continue the titration until the color of the methyl orange is destroyed. Make a blank titration using exactly the same volume of hydrochloric acid, water, and 2 drops of methyl orange indicator solution, replacing the volume of arsenite or antimony test solution by additional distilled water. The blank should run about 0.12 to 0.14 ml. RESULTS
I n order to determine the stability of the standard sodium hypochlorite solution prepared in the manner directed above, titrations were performed as detailed a t intervals during 3 years. I n all, five series of experiments were run. Titrations were performed in triplicate and the results averaged. The results obtained in three of these series are representative. I n the first the initial normality was 0.004970; after 102 days it waa still 0.004970; the maximum variation within this period was $0.000008. I n the second series, initial normality was 0.005572; after 175 days this had degraded to 0.005543, the maximum varia-
stability at ~ I H 1: 206
ANALYTICAL EDITION
March, 1944
tion within the period being -0.000021. I n another series, the initial normality was 0.00500 and at the end of 56 days was 0.004977 with a maximum variation of -0.000020. The other two series gave comparable results. As a check on the accuracy of sodium hypochlorite solution as a titrimetric reagent, arsenic and antimony were determined in standard solutions, using both the potassium bromate and the sodium hypochlorite methods. DISCUSSION
I n the determination of microquantities of antimony, the potassium bromate method (1) was not entirely satisfactory for a number of reasons, the principal ones being that titrations had to be performed almost a t the boiling point in strong hydrochloric acid solution with the consequent production of relatively copious and irritating fumes of hydrogen chloride and that the blank, using methyl orange as the indicator, was large. The blank obtained using some of the indicators suggested by Smith and Bliss (11) was even larger. Standard sodium hypochlorite solution has several marked advantages. It is an economical reagent. One can perform direct titrations with it. It is unnecessary to perform the titrations at elevated temperatures, eliminating any danger from the irritating fumes of hydrochloric acid. The blank is smaller than that obtained with potassium bromate titrations. Hypochlorite titrations can be performed under conditions of low acid concentration without apparent decrease in accuracy. From 0.1 to 1mg. of antimony or arsenic per 10 ml. of sample solution can easily be estimated. Titrations can be made in glass-stoppered bottles,
207
glass-stoppered Erlenmeyer flasks, or iodine flasks, if desired, to minimize losses attributed to volatility. Many of the indicators mentioned by Smith and Bliss, and Kolthoff and Stenger can be used instead of methyl orange without increase of the blank. The precision compares favorably with that of other methods, as can be seen from the reproducibility of results. Several precautions must, however, be observed in using sodium hypochlorite solution as a titrimetric reagent. It must be preserved in brown, glass-stoppered bottles. It may be kept at room temperature without deterioration over considerable periods of time. Keeping the solution at lower temperatures is perhaps preferable. The optimum condition3 for the titrations are a volume of at least 35 to 40 ml. with an acid concentration equivalent to 5 ml. of concentrated hydrochloric acid. LITERATURE CITED
Anderson, ISD. ENQ.CHEM.,A s . 4 ~ ED., . 11, 224 (1939). Chapin, J . Am. Chem. Soc., 56, 2211 (1934). DenigBs, J . pharm. chim., [5] 23, 101 (1891). Jackson and Parsons, IXD. ENG.CHEM.,ANAL.ED.,9, 14 (1937). Jellinek and Kresteff, 2. anorg. Chem., 137, 333 (1924). Jellinek and Kuhn, Ibid., 138, 81 (1924). Kolthoff and Stenger, IND.ENQ. CHEM.,ANAL. ED., 7, 79 (1935).
Koppeschaar, 2. a n d . Chem., 15, 233 (1876). Lunge-Berl, “Chemisoh-technische Untersuchungsmethoden”, Vol. 1, 7th ed., Berlin, J. Springer, 1921. Manchot and Oberhauser, Z . anorg. Chem., 130, 161 (1923). Smith and BlisJ, J . Am. Chem. SOC.,53, 2091 (1931).
Thiosulfate Washers in Alkoxy Microdeterminations E.
P. WHITE, Chemical Laboratory,
Animal Research Division, Department of Agriculture Wellington, N e w Zealand
Determinations of methoxy and methylimide groups in which thiosulfate alone is used as a washer give values considerably lower than theoretical. This is due to the solubility of the methyl iodide in the washer, and a subsequent reaction. Ethoxy and ethylimide determinations are not subject to this loss. The effect of thiosulfate can be eliminated b y using as washer thiosulfate dissolved in saturated sodium chloride, or b y adding cadmium sulfate as in the standard gravimetric procedure. The minor errors in determinations using as washers water, phosphorus suspension, or 0.570 sodium carbonate, are insignificant in comparison with that due to thiosulfate. A rapid distinction between ethoxy and methoxy can be made b y doing a determination with a good washer, such as 0.5% carbonate, then with 5% thiosulfate) the methoxy value will be reduced to 55 to 70% of the original, while ethoxy remains unchanged.
W
ORK on alkaloids in this laboratory required the development of micromethods of alkoxy and akimide
determination. The apparatus used was that of Pregl for alkoxy and that of Friedrich for alkoxy and alkimide determinations, and the procedure was essentially that of modern textbooks of microchemistry. Three to 5 mg. of material were weighed on tinfoil, dissolved in phenol and acetic anhydride, heated with hydriodic acid, and passed through a washer of 5% thiosulfate containing 0.5% sodium carbonate. Final estimation was by the Viebock-Brecher method. Preliminary experiments with the Friedrich apparatus showed that the values obtained with vanillin and several alkaloids did not agree with theory, calculation being from first principles. The titration obtained with all methoxy-containing substances was only 50 to 70% of the theoretical, while ethoxy values agreed closely with theory. The method was then examined in detail and many
of the more obvious possible sources of error were eliminated. The same effect was found in the Pregl apparatus. The only factor not eliminated appeared to be the washing solution. Consequently, red phosphorus suspension, the original washer of Pregl ( 9 ) , as well as water and 0.5% sodium carbonate was tried. These gave theoretical results in the Pregl apparatus, and values some 5% low in the Friedrich apparatus. The use of thiosulfate as a washer was then investigated, and the literature searched for counterindications to its use. Thiosulfate washers are almost universally used and recommended by the later workers in microchemistry. EXPERIMENTAL
Known amounts of methyl iodide were introduced into the Friedrich apparatus without any hydriodic acid, drawn through various washers, and titrated In the ordinary way. With no washer, or with water, phosphorus suspension, or 0,5y0carbonate the recovery was almost theoretical. With thiosulfate there was only 50 to 65% recovery, thus confirming the effect of thiosulfate. A survey of results obtained with various washers in the Pregl apparatus is given for vanillin in Table I, and for phenacetin in Table 11, which show that the effect of thiosulfate on the methoxy value is detected when 1 mi. of a 1.57, solution is used. This effect becomes much larger when 5 to 10% thiosulfate is used, while with very high concentrations (40%) theoretical values are again obtained. With ethoxy there was no detectable effect in any concentration. DISCUSSION
The effect of thiosulfate is explained as a result of two factors: (1)the solubility of the alkyl halide in the washing solution, with
