Determination of Available Chlorine in Hypochlorite Solutions by Direct Titration with Sodium Thiosulfate VIRGILA. WILLSON,Montana Livestock Sanitary Board, Helena, Mont.
T
H E use of sodium thiosulfate for the direct volumetric determination of available chlorine in a hypochlorite solution has been considered unsatisfactory and the literature is confusing as to the exact reaction which takes place. In many discussions of the subject, the author avoids the reaction of chlorine by giving the iodine reaction and leaves the impression that chlorine acts in the same manner as iodine and according to the following well-known equation: 2NazSzOs Clz +NazSIOe 2NaCl (1)
The strength of the arsenite solution was so adjusted that it reacted with equivalent quantities of a 0.1 N iodine solution when using starch as an indicator. The reaction of sodium thiosulfate on a hypochlorite solution was determined (a) by titrating an acetic acid solution of sodium hypochlorite with a 0.1 N sodium thiosulfate solution and using starch-potassium iodide paper as an outside indicator; (b) by adding an excess of potassium iodide to an acetic acid solution of sodium hypochlorite, and titrating the liberated iodine with the same thiosulfate solution used in (a), while using starch solution as an internal indicator. Mohr (4) states that eight equivalents of chlorine from Fifty cubic centimeters of chlorine solution required in (a) chlorine water are consumed in the complete oxidation of 2.88 cc. and in (b) 23.05 cc. of sodium thiosulfate. Therefore, one mole of sodium thiosulfate, but infers that constant 23.05/2.88 equals 8.003 equivalents of chlorine reduced per results cannot be obtained when this method is used for de- equivalent of iodine reduced by sodium thiosulfate from an termining available chlorine. Lunge (3) gives the theoretical acetic acid solution of a hypochlorite. reaction for chlorine on sodium thiosulfate as follows: Several hundred analyses of solutions of sodium and calcium hypochlorite, with chlorine concentrations varying from 25 Na2SzOa 8C1 5Hz0---f HzS04 NapSO4 8HC1 (2) parts per million to 1.25 per cent of available chlorine, have He states that this reaction does not go to completion in the been made by the use of this method and the results have manner indicated, and that in all cases the amount of thio- been found to check within an error of hO.2 per cent of sulfate consumed is greater than is required by Equation 2 the available chlorine as determined by the arsenic trioxide and less than is required by Equation 1. He suggests the method. Tests have been made a t temperatures from 10' possibility of two or more reactions taking place to form to 35" C. and with concentrations of acetic acid from barely different oxidation products of sulfur, and states that sulfur, acidic to 33.3 per cent glacial acetic acid. At high temperasulfur dioxide, and hydrogen sulfide were all observed during tures and with a large excess of acetic acid, there is a loss of different conditions of the experiment. Di6nert and Wanden- chlorine before titration, but the reaction goes t o completion bulcke (a), when attempting to formulate a method for to form the sulfate, as has been shown by checking the concalculating the amount of sodium thiosulfate solution re- centration of available chlorine with a standard arsenic quired to neutralize the chlorine in chlorine water, ob- solution. Tables I, 11, and I11 are submitted to show the effect of tained results which were erratic and which were influenced by the alkali and carbonate content of the solution. Bloxam acetic acid, strong alkalies, and high temperatures, respec(1) gives Reaction 2 without discussion. Treadwell and tively, on the reaction. Hall (6) state that chlorine and bromine react like iodine on TABLE I. EFFECT OF ACETICACID sodium thiosulfate when they are not present in excess, but that when they are present in excess the tetrathionate is Reaction temperature 25O C. Indicator, starch-pot assium iodide paper. sohtion, 20 cc. equals 19.16 CC. of 0.1 N oxidized to the sulfate and sulfuric acid with the deposition &oncentration of hypdchlorite arsenite solution.) of sulfur which in turn is gradually oxidized to sulfate. 10 PERCENT ACETICACID 0.1 N REACTION T O LITMUB In an attempt made in this laboratory to make use of either HYPOOHLORITE SOLUTION SOLUTION SODIUM Before After UBED A D D ~ D THIOSULFATE titration titration Reaction 1 or 2 in the volumetric estimation of available cc. cc. cc. chlorine in a hypochlorite solution, using starch-potassium Basic Basic 0 36.30 20 Basic Basic iodide paper as an outside indicator, the results obtained were 0.1 38.03 20 Basic Basic 39.60 0.2 20 erratic and the amount of thiosulfate consumed was not Basic 42.00 Basic 0.3 20 Basic Basic 45.05 0 . 4 20 correct for the satisfaction of either equation. When, howBasic Basic 49.20 0.5 20 Basic Basic ever, the hypochlorite solution was acidified with acetic 63.60 20 0.6 Acidic Acidic 19.12 0.7 20 acid before titration, the results were found to be different Acidic Acidic 0.8 19.14 20 Acidic Acidic 19.13 0.9 20 and indicated the completion of the reaction as shown by Equation 2, where eight equivalents of chlorine are conTABLE11. EFFECT OF STRONQ ALKALIES sumed in the complete oxidation of one mole of sodium (Reaction temperature 25O C. Indicator, starch-potassium iodide paper. thiosulfate. No record was available to show that this re- Concentration of hypdchlorite so!ution, 20 cc. equala 18.40 co. of 0.1 N aclion had ever been tested out in an acetic acid solution. arsenite solution.)
+
+
+
+
+
+
O F HYPOOHLORITE HYPOCHLORITETREATMDNT SOLUTION SOLUTION BEFORE TITRATION Cn . - .. 100. of 1%NaOH added 20 10 cc. of 1 NaOHadded 20 1 cc. of 5 % NaaCOs added 20 10 cc. of 6 Na2COa added 20 1 cc. of 5 7 NaHCOa added 20 10 00. of 5 % NaHCOs added. 20
EXPERIMENTAL DATA PREPARAT~ON OF SOLUTIONS. The strength of the sodium thiosulfate solution was so adjusted that 50 cc. of the thiosulfate reacted with 6.25 cc. of a 0.1 N iodine solution when using starch as an indicator. 44
0.1 N S O D I U u THIOSULFATE
cc.
34.25 26.30 35.10 35.00 42.40
56.50
January 15, 1935
ANALYTICAL EDITION
TABLE 111. EFFECT OF HIGHTEMPERATURES (Indicator, starch-potassium iodide paper.) 10 PERCENT 0.1 N ACSTICACID SODIUM
HYPOCHLORITE SOLUTION
* c.
cc.
ADDBD
0.1 N
SODIUM THIOSULFATE -4RSnNITE
CC.
CC.
CC.
DISCUSSION OF RESULTS The reaction of sodium thiosulfate on a hypochlorite is influenced by the alkalinity of the solution, as is shown by Tables I and 11. As the alkalinity is decreased, the amount of thiosulfate required for the destruction of the chlorine is increased until the solution is made acidic, when a different reaction takes place and goes to completion to form the sulfate according to Equation 2. Di6nert and Wandenbulcke (2) state that the alkalinity of a hypochlorite solution after titration with sodium thiosulfate is the same as the alkalinity of the hypochlorite solution determined after the destruction of the chlorine with hydrogen peroxide. This is not in conformity with the results obtained in this laboratory and the following explanation concerning the mechanism of the reaction is offered: When Reaction 1 takes place with a hypochlorite solution it is correctly written as:
+
2Na2S20s NaOCl
+ HzO+NazSIOa+ 2NaOH + NaCl (3)
where one equivalent of alkali is liberated per mole of sodium thiosulfate consumed. Under the same conditions Reaction 2 should be written:
+
Na2S20, 4NaOCl
+ HzO-+
4NaC1
+ Na2S0, + H2SOc (4)
where two equivalents of acid are liberated per mole of sodium thiosulfate consumed. Adding Equations 3 and 4 we get Equation 5 which is a neutral reaction. 3Na&3SO8
+ 5NaOCl+
Na&&Oe
+ 5NaCl + 2NazSOa (5)
From these reactions we may conclude: 1. If Reaction 3 takes place in excess of Reaction 4, the
alkalinity of the hypochlorite solution will be increased. 2. If Reaction 4 takes place in excess of Reaction 3, the alkalinity of the hypochlorite solution will be decreased. 3. If Reactions 3 and 4 take place to form Reaction 5, the alkalinity of the solution will not be affected. In all cases of titrating hypochlorite solutions which contain sodium hydroxide or carbonate, with sodium thiosulfate, the amount of thiosulfate consumed indicates that Reaction 4,although not quantitative, is the predominating one. With hypochlorite solutions that contain an excess of bicarbonate or are neutral in reaction, the thiosulfate consumed exceeds the theoretical amount required for Equation 5, and an alkaline reaction is produced. This accounts for the fact that when a neutral or an alkaline solution of a hypochlorite is reduced with sodium thiosulfate, the resulting soh tion is always alkaline, and explains why in some cases the alkalinity after titration with sodium thiosulfate is the same as the alkalinity of the solution after the chlorine has been destroyed with hydrogen peroxide. In no case was the precipitation of sulfur observed or sulfur dioxide or hydrogen sulfide liberated when sodium thiosulfate was added to either an acetic acid or an alkaline solution of a hypochlorite. The experiments as here conducted have been with solutions of alkaline hypochlorites and with solutions of hypochlorites containing a small excess of acetic acid and consequently hypochlorous acid, while chlorine
45
solutions as previously referred to may have in some instances implied solutions of chlorine in water which contained hydrochloric acid in equivalent proportions to the hypochlorous acid, In this case the precipitation of sulfur and the formation of sulfur dioxide are to be expected. Very dilute hydrochloric or sulfuric acids act like acetic acid, but with these acids there is danger of a loss of free chlorine and also of the formation of sulfur and sulfur dioxide. Chlorates which may be present in hypochlorites do not interfere with the reaction of sodium thiosulfate in an acetic acid solution of a hypochlorite and this method gives a measure of the available chlorine only (5). In processes where sodium thiosulfate is used as an antichlor, it is believed that the amount of thiosulfate required for the destruction of the chlorine can be calculated, and that a great saving can be made in the amount of thiosulfate by properly controlling the hydrogen-ion content of the solution and forcing the antichlor reaction to completion to reduce eight equivalents of chlorine per mole of sodium thiosulfate. METHODOF ANALYSIS REAGENTS REQUJRED.0.1 N sodium thiosulfate, starchpotassium iodide paper, and dilute acetic acid (10 per cent glacial acetic). PREPARATION OF 0.1 N SODIUM THIOSULFATE SOLUTION. This solution is prepared by dissolving slightly over 3.1025 grams of the crystallized salt per liter of solution and adjusting the concentration so that it reacts with equivalent volumes of an iodine solution containing 1.5865 grams of pure iodine per liter of solution. The solution is then 0.1 N to chlorine when used for titrating an acetic acid solution of a hypochlorite and is 0.0125 N t,o iodine. METHOD. The hypochlorite is diluted with water to contain approximately 0.1 per cent of available chlorine. Fifty cubic centimeters of this solution are pipetted into an Erlenmeyer flask and dilute acetic acid is added until a drop of the solution, when taken out on a glass rod, shows an acid reaction t o litmus before the litmus is completely bleached. The thiosulfate solution is then run into the hypochlorite solution from a buret, a few cubic centimeters at a time, until a drop of the reacting solution fails to give a chlorine reaction with starch-potassium iodide paper. A second or third titration may be necessary in order to obtain the exact end point. The concentration of available chlorine in the original hypochlorite is then calculated from the amount of dilution and the volume of sodium thiosulfate used. SUMMARY Experiments conducted in this laboratory have demonstrated that sodium thiosulfate, when added to a dilute acetic acid solution of a hypochlorite, is completely oxidized to the sulfate and that eight equivalents of chlorine are used per mole of sodium thiosulfate. The reaction can be used as a quantitative method for determining available chlorine in a hypochlorite. The sodium thiosulfate solution is standardized by titrating against an iodine solution of known concentration. The end point of the reaction is determined by the use of starch-potassium iodide paper as an external indicator. When working a t laboratory temperatures of from 20” to 25” C., this method gives results which check with those obtained when using a standard arsenite solution for reducing the chlorine. LITERATURE CITED (1) Bloxam, A. G., “Inorganic and Organic Chemistry,” 11th ed., p. 367, Philadelphia, P. Blakiston’s Son & Co., 1923.
(2) DiBnert, F., and Wandenbulcke, F., Compt. rend., 169, 29, 30 (1919). (3) Lunge, G., Ber., 12,404 (1879). (4) Mohr, C. F., “Chemisoh-Analytischen Titrimethode,” 2nd ed., p. 252, Braunsohweig, F. Vieweg und Sohn, 1862. (5) Sutton. F., “Volumetric Analysis,” 11th ed., p. 186, Philadelphia, P. Blakiston’s Son & Co., 1924. (6) Treadwell, E”. P., and Hall, W. T., “Analytical Chemistry,” 6th ed., Vol. 1, p. 325, New York. John Wiley & Sons, 1927.
RECEIVED October 16, 1934
