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 Title Page
 Table of Contents
 Introduction and review of the...
 Statement of problem
 Experimental methods
 Experimental data
 Discussion of results
 Summary
 Bibliography
 Acknowledgement
 Biographical items
 Copyright














Title: oxidation of sulfides by chlorine in dilute solutions
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Table of Contents
    Title Page
        Page i
    Table of Contents
        Page ii
    Introduction and review of the literature
        Page 1
        Page 2
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    Statement of problem
        Page 13
    Experimental methods
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    Experimental data
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    Discussion of results
        Page 101
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    Summary
        Page 108
        Page 109
        Page 110
    Bibliography
        Page 111
        Page 112
        Page 113
        Page 114
    Acknowledgement
        Page 115
    Biographical items
        Page 116
        Page 117
    Copyright
        Copyright
Full Text
THE OXIDATION OF SULFIDES BY CHLORINE IN DILUTE AQUEOUS SOLUTIONS
By
JAMES BROWN GOODSON, JR.
A DISSERTATION PRESENTED TO THE GRADUATE COUNCIL OF THE UNIVERSITY OF FLORIDA IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF DOCTOR OF PHILOSOPHY
UNIVERSITY OF FLORIDA July, 1950


TABLE OF CONTENTS
Page
INTRODUCTION AND REVIEW OF THE LITERATURE .............. 1
STATEMENT OF PROBLEM................................... 13
EXPERIMENTAL METHODS................................... 14
Preparation and Standardisation of Solutions ....... 14
Description of Apparatus *..*.**....*.*. 23 Experimental Procedures ..*. 30
EXPERIMENTAL DATA *...*...*......**...... 42
Precision of Determinations ...........* 44
Effect of Concentration on the Reaction ............ 47
Effect of Time on the Reaction ..................... 55
Effect of Temperature on the Reaction ........ 59
Effect of Hydrogen-ion Concentration on the Reaction 67 Effect of Ionic Strength on the Reaction ........... 83
Effect of Chloride Concentration on the Reaction ... 94 DISCUSSION OF RESULTS .................................. 101
SUMMARY .......... ...................................... 108
BIBLIOGRAPHY........................................... Ill
ACKNOWLEDGEMENTS.......................................115
BIOGRAPHICAL ITEMS ....... *............................. 110
COMMI TTJBfi REPORT ................ . ..................... 117


INTRODUCTION AND REVIEW OP THE LITERATURE
Sulfides are commonly found in the natural waters of several regions of our land, where their occurrence is noted with far greater frequency in ground waters, which are generally devoid of oxygen, than in surface waters. The state of Florida lies in a region where sulfide concentrations in natural waters up to several parts per million ore not uncommon. The origin of these sulfides is attributed to biological and chemical processes whereby sulfur compounds are decomposed and sulfur In a free or oxidized state Is reduced.
Experience has taught us that the presence of sulfides In a water renders that water highly undesirable for domestic and industrial usages. There have been numerous instances where the presence of sulfides In an industrial or domestic water supply has been found to be responsible for offensive odors, excessive corrosion and conditions leading to heavy growths of micro-organisms with the undesirable consequences attached thereto. Furthermore, sulfides have been found to be responsible for extensive damage to greensand zeolite beds. The result of such observations has been the realization of the urgent need for the removal of sulfides from waters to be used for domestic and industrial purposes.
It has been common practice for many years to effect the removal of sulfides from water supplies by aeration. By this
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method part of the sulfide, the amount being largely dependent upon the hydrogen-ion concentration of the water, is expelled to the atmosphere as hydrogen sulfide, and the remainder is oxidized to some degree by the dissolved oxygen in the aerated water* the method is attended by many difficulties of an engineering nature, and its efficiency is dependent upon several chemical and physical variables* Therefore, the over-all effectiveness of the method may be seen to vary considerably from one installation to another.
Chlorinetlon as a means for sulfide removal from natural waters is of rather recent origin. It appears to date from 1926, (1) in which year it was resorted to at Beverly Hills, California, after aeration alone had been proven to be ineffective for complete sulfide removal. As late as 1936 Cox, (2) in reviewing progress in the elimination of tastes and odors from water supplies, expressed interest in the apparent novelty of hydrogen sulfide removal by chlorination when he made the following statements "Mention should be made, however, to the practice at Holland, Hew York, where aeration is used to remove hydrogen sulfide from well water. The Holland supply is also disinfected with chlorine, and it is interesting to note that the chlorine will react with any residual hydrogen sulfide present In the aerated water and completely remove It.*. Earlier mention is made of the effect of hypochlorite on hydrogen sulfide in septic sewage liquors. Rldeal, (3) in reporting the results of a long


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series of experiments dealing with hypochlorite treatment of septic sewage liquors in the 1908 report of the Royal Cora-mission on sewage, calls attention to the fact that 35 to 50$ treatment with hypochlorite solution is sufficient to do away with the offensive hydrogen sulfide smell leaving in its place a comparatively inoffensive odor of spent bleaeh and fresh sewage*.
Discrepancies in the literature in regard to the dosage of chlorine required for sulfide removal indicate the need for fundamental research concerning the reaction between sulfides and chlorine in dilute aqueous solutions. In the usual case in the literature it is assumed that chlorine oxidizes the sulfide to free sulfur, and the dosages required are based on the stoichiometric relationships pertaining to this reaction. Hoover (4) states that the reaction between chlorine and hydrogen sulfide is as follows! RgS / Clg Z 2HC1 / S. He concludes that for the removal of each pound of hydrogen sulfide 2.1 pounds of chlorine are required. Pomeroy and Bowlus (5) have the following to sayt *fhen chlorine is added to a pure water solution of sulfides, it requires only 2.22 parts of chlorine to destroy one part of sulfide as indicated by the following reaction; Gig / HgS S S / 2HC1.*. Thus it is seen that these authors are basing their dosages on the stoichiometric relationships involved in the reactions they have written. However, Pomeroy and Bowlus (5) in determining th chlorine dosages just sufficient


4
to eliminate the sulfides from various samples of septic sewage found that the ratio of chlorine to sulfides varied from 3 to 9, averaging 5,3. They explain the higher dosages in the case of sewage by pointing out that other reducing agents would consume a part of the chlorine. Experiments by Powell and von Lossberg (6) indicate that the chlorine dosage for removal of hydrogen sulfide from some natural waters in which they were interested approximates the stoichiometric value dictated by the following reaction! 4C1^ / 4Hg0 /
V = H2S04 / SHC1- stoichiometric value calculated from the relationships involved In this reaction is seen to be 8.84 parts of chlorine to one part of hydrogen sulfide expressed as sulfide. In actual practice chlorine dosages for sulfide removal are generally determined by chlorine demand tests (7) with the particular water in question. However, Wallace and Tiernan Co., Inc., has patented a method (8) whereby the amount of chlorine to be added Is indicated by means of a change in an electrode potential as chlorine Is added to the aqueous material. These methods require no knowledge of the reactions involved in the oxidation of the sulfides.
The oxidation of sulfides in aqueous solutions has been studied by many investigators using a variety of oxidizing agents* A review of the work of several of these investigators indicates that the products of the oxidation vary with the relative oxidising potentials of the oxidising


couples used* In the following discussion couples with oxidation potentials In the vicinity of or greater than one volt are considered to he strong oxidants, and those with potentials less than one volt are considered to he weak oxidants* Among the weakest oxidants studied were neutral to alkaline solutions or suspensions of potassium chromate (9,10), potassium diehromate (11), ammonium chromate (12), lead chromate (13), silver chromate (13), mercurous chromate (14), barium chromate (15) and solutions of various water soluble aromatic nitr compounds (16)* The principal pro-duets of the oxidation of alkali sulfides and hydrogen sulfide, by these oxidants were found to be polysulfides, free sulfur, thlosulfates, sulfites and sulfates* The formation of sulfites and sulfates appears to be dependent upon the temperature and alkalinity of the solutions, higher temperatures and lower hydroxyl-ion concentrations favoring the production of sulfates. Bullock and Forbes (16) show that the oxidation of sulfides by such mild oxidants as aromatic nitro compounds progresses only as far as free sulfur. They attribute the presence of the other final products, thlosulfates and polysulfides, to a secondary reaction between "active* sulfur and hydroxyl-ion, which they write as followsi 60H / 12S SgOg""* / 2Sg / 3Hg0. The investigators-describe "active" sulfur as that sulfur set free by an oxidation of a sulfide or in some similar way. This work of Bullock and Forbes brings to mind the oxidation of hydro-


. 6 -
gen sulfide by Iodine, which Is considered to be a weak oxidant. That oxidation proceeds quantitatively to free sulfur in an acid medium (1?) and Is the basis for a well known method for the quantitative determination of hydrogen sulfide in aqueous solutions. However. In alkaline solutions sulfate appears as a product of the oxidation (17, 18).
Some moderately weak oxidants that have been used in studies of the oxidation of alkali and hydrogen sulfides are oxygen in neutral to slightly alkaline solutions (19, 20), nitrous acid (21) and solutions of calcium permanganate (22), silver permanganate (22), ammonium permanganate (23), barium permanganate (23) and chromium trloxide (24). The principal products resulting from the oxidations with these oxidants are free sulfur, thiosulfates, sulfites and sulfates. It is noted that in these studies polysulfides are not mentioned among the products found.
Strong oxidants that have been employed Include oxygen in acid solutions (20), potassium iodate in acid solutions (25), nitric acid (26), paraperiodic acid (27) and a neutral solution of potassium permanganate (28). Free sulfur and sulfates were found to be the only end-products with these oxidants except in the case of the oxidation with neutral potassium permanganate, in which ease it was found that under certain conditions relative to concentrations a dlthlonate occurs as a final product.
The preceding review of the work of some of the invest!-


- 7 -
gators who hare studied the oxidation of sulfides in aqueous solutions indicates that the hydrogen-ion concentration of the medium in which the oxidation takes place is a very Important factor in determining the characteristics of the reaction. Free sulfur is apparently the primary oxidation product of sulfides and may he the only end-product In acid solutions when weak oxidants are employed (16, 17). In basic solutions there are changes noted in the final products of the oxidation (17, 18, 20). These changes are evidently explained by effects of the hydrogen-Ion concentration on the oxidant, on the sulfide equilibria relationships and on the Intermediate products of the oxidation*
The oxidation potentials of the large majority of the oxidants used by these investigators are dependent upon the hydrogen-ion concentration of the medium (29). For example, it is noted that the standard oxidation potential for the water-oxygen couple in acid solutions Is -1.229 volts, which Indicates that oxygen is a strong oxidant in acid media. The potential for the same couple in a neutral solution Is only -0.815 volt, indicating a moderately weak oxidant. Another possible effeet of the hydrogen-ion concentration on certain oxidants is illustrated in the specific case where nitrous acid was employed as the oxidising agent.
Mention was made by Bagster (21) that the results of the study suggest that free nitrous acid is the active agent rather than nitrite ion. If this is true the importance of


the hydrogen-Ion concentration lies In Its influence on the equilibrium involving free nitrous acid and nitrite ion in the reaction mixture. Still another effect of the hydrogen-ion concentration on an oxidant which leads to subsequent changes in oxidising characteristics is noted In the case of iodine (17). Although the oxidation potential of the iodide-iodine couple Is not dependent upon the pH of the solution at pH values below 8, at higher pH values the iodine will react with the hydroxyl-ion to yield hypoiodite (30) thereby altering the oxidising properties of the mixture. It is inferred from the work of Kapp (31), who deduced from experiments dealing with the oxidation of alkali sulfides that hydrogen sulfide should be more easily oxidised in aqueous solutions than sodium sulfide, that there might be a possible effect of hydrogen-ion concentration on the oxidation of sulfides from the standpoint of its Influence on the sulfide equilibria relationships. The influence of the hydrogen-ion concentration on Intermediate oxidation products of sulfides may be exemplified by the reactions between sulfur and alkali and alkaline earth hydroxides in aqueous solutions. Tartar (32) found that the primary reaction is as follows* 60H~ / 8S s 2Sg-~ / SgOg""" / 3HgO. An excess of sulfur was found to yield pentasulflde by a secondary reaction. As pointed out by Bullock and Forbes (16) the reactions do not proceed rapidly at 25 C. with ordinary rhombic sulfur, but when the sulfur is in the "active* state the reactions are quite rapid


- 9 -
even at this temperature* A further example of the effect of hydrogen-Ion concentration on a product of the oxidation of sulfides is seen in the stability exhibited by thiosul-fates in alkaline solutions (33, 34)*
How let us consider some of the properties of chlorine as an oxidising agent in aqueous solutions* Latimer (29) gives -1.3583 volts as the best value for the standard oxidation potential of the chloride-chlorine couple, but he points out that the potential becomes meaningless in alkaline solutions because of the hydrolysis of chlorine and the formation of hypochlorite. He lists the standard oxidation potential of the ehlorlde-hypochlorlte couple as *0.94 volt. From this latter value it is calculated that the approximate values of the oxidation potential of chlorine in aqueous solutions having pH values of 9 and 7 are -1.24 and -1.35 volts, respectively. Thus it is seen that chlorine is a strong oxidant throughout the normal pH range of natural waters. Biggins (35, 36), Hideal and Evans (37) and Remington and Trimble (38) have studied the effect of acids on the oxidising properties of hypochlorite solutions. It is apparent from the works of these investigators that the oxidising power of hypochlorite solutions can be markedly increased by the addition of weak acids. Rldeal (39) suggests that free hypochlorous acid is the effective oxidiser In such solutions and attributes the effect of the added acids to the resulting Increase In concentration of that com-


- 10 -
ponent of the oxidizing mixture, Rideal and Evans (37) call attention to the values of the dissociation constants of hypoehlorous and carbonic acids to shov that free carbonic acid will liberate free hypoehlorous acid from hypochlorite solutions* Biggins (38) points out that whereas the addition of an excess of boric acid to a hypochlorite solution results in a solution of very energetic bleaching properties, the addition of an excess of hydrochloric acid gives a solution of very weak bleaching properties. Be also attributes the oxidising properties of such solutions to the active mass of free hypoehlorous acid and explains that the boric acid liberates free hypoehlorous acid while the hydrochloric acid liberates free chlorine* Weiss (40) is discussing the kinetics of chlorine bleaching claims that the active agent is chlorine monoxide or undlssoelated hypoehlorous acid. He mentions another interesting point in connection with the oxidising activities of chlorine solutions, and that is that there Is an apparent maximal rate of attack on cellulose fibers at a certain pH value, the rate falling rapidly with either an increase or decrease in pH, Blakely (41) measured the oxidation potentials of hypochlorite solutions having pH values from 2 to 13 and found that a maximum is indicated at pH 7,0, It was discovered by Hlggins (42) that chlorides have an accelerating effect en the bleaching action of chlorine solutions* The effect was found to be an immediate one, after which action the solutions behave as though the chlo-


rides were not present. Chlorides produced hy the reduction of the hypochlorites during the bleaching reaction appear to have a negligible accelerating effect. It is mentioned by Higglns (43) that there Is a secondary reaction between hypoehlorous acid and neutral chloride whereby nascent chlorine of energetic bleaching properties is produced.
Information in the literature in regard to the oxidation of sulfide solutions by chlorine appears to be rather meager. Stock (44) reports that the oxidation of a dilute hydrogen sulfide solution in anhydrous, liquid hydrogen chloride yields sulfur as the oxidation product, which lends support to the view that sulfur is the primary product in the oxidation of such sulfides, Perel'man and Lelyakina (45) found that hydrogen sulfide in acetylene gas is quantitatively oxidised to sulfate when passed through a solution of sodium hypochlorite and propose a quantitative method for the determination of hydrogen sulfide in acetylene based en this reaction. Some comprehensive work has been done by Choppin and Faulkenberry (46) on the oxidation of aqueous sulfide solutions by hypochlorites, these investigators performed their studies using reaction solutions that were considerably more concentrated than the solutions encountered in water works practice, the solutions varying from 40 to 2000 parts per million in sulfide concentration. They established the fact that the end-products of the oxidation are sulfur and sulfate, the ratio depending upon such factors as relative concentrations of the original reaetants, hydrogen-ion concentration of the reaction medium.


12 -
temperature, standing time and rate of addition of reaetants. The stand is taken that sulfur is the primary product of the oxidation, whereas sulfate results as the end-product of secondary reactions that may occur simultaneously with the primary reaction. The effect of the above-mentioned factors on the ratio of sulfate to sulfur produced is attributed to the influence of these factors on the secondary reactions. For example, higher temperatures were found to increase the proportion of sulfate, and pH values of 13.8 or higher were found to result in a quantitative oxidation to sulfate. The result in each of these cases is explained by the investigators on the basis of the equation, 60H" / 8S z ZS^ / SgOg** / 3HgO, where increased temperatures and high a Ileal inl ties favor the solution of sulfur (32, 16). The products are subsequently oxidised to sulfates (47). It was also noted that there is a quantitative oxidation of sulfide to sulfate at pH values of 2 or less. The explanation given for this depends upon the presence of chlorine monoxide in acid solutions of hypochlorites (48) and its function as a reagent for the re-solution of colloidal sulfur. Between pB values 2 and 13.8 the proportion of sulfate was found to decrease to a minimum at a pH value in the vicinity of 10.


STATEMENT OF PROBLEM
The object of this investigation was to make a study of the oxidation of sulfides in very dilute aqueous solutions, sueh as those normally encountered in water works practice* Particular attention was devoted to the effect on the reaction of hydrogen-ion concentration, time, concentrations of react-ants, temperature, ionic strength and chloride concentration* Since the study was made with a view toward application in the water treatment field, the limits adopted for the various variables have corresponded where possible to those commonly experienced in the water works field*
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EXPERIMENTAL METHODS
The methods chosen for use in these investigations were based upon the familiar chlorine demand method (7). Sulfide Solutions of various known characteristics were accurately-made up In the absence of oxygen in a series of reaction vessels, known dosages of chlorine were added, and the amount of chlorine or sulfide remaining after a measured elapse of time under controlled conditions was accurately determined by lodometrie or iodimetrle methods, respectively* From the results of the residual chlorine or sulfide determinations it is possible to calculate the values of the ratio of chlorine reacted to sulfide reacted. The details of the procedures employed are given later*
/.
Preparation and Standardization of Solutions
Standard 0.1 J! potassium dichromate solution. Some reagent grade potassium dichromate was pulverised in an agate mortar and dried in an oven at 150 200 C. 9.308 grams of the material were dissolved in distilled water and diluted exactly to 2 liters.
Standard 0.01 N_ potassium dichromate solution. 200 ml, of the standard 0.1 N solution were transferred to a 2 liter volumetric flask by means of a calibrated 100 ml. volumetric pipet, and then the solution was diluted exactly to 2 liters.
- 14 -


15 -
Sulfuric acid. Concentrated reagent grade sulfuric acid vas used where this acid was called for in the procedures.
Potassium iodide. Reagent grade potassium iodide crystals that hare been tested for the absence of lodate were employed in the investigation.
Starch indicator solution. 5 grams of potato starch were mixed with a little cold water in a mortar and ground to a thin paste. The mixture was poured into a liter of boiling distilled water, stirred and allowed to settle overnight. The clear supernatant was used as the indicator solution. Since the solution is subject to biological decomposition, salicylic acid (1.25 grams per liter) was added as a preservative*
Zinc acetate solution. 240 grams of reagent grade zinc acetate were dissolved in one liter of distilled water.
0.01 |J Iodine solution. 2*54 grams of reagent grade iodine were dissolved in several ml. of water containing S grams of lodate-free potassium iodide. This solution was diluted to approximately 2 liters* No standardization was necessary.
Hydrochloric acid. Concentrated reagent grade hydrochloric acid was used where this acid was called for In the procedure.
Acetic acid. 500 ml. of glacial acetic acid was diluted to one liter with distilled water.


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0.1 ] sodium thlosulfate solution. Approximately 206 grams of C. P. sodium thlosulfate pentahydrate were dissolved in 8 liters of recently boiled, cooled distilled water containing 0,8 gram of sodium carbonate. A few ml. of chloroform were added as a preservative, and the solution was allowed to stand for several days before standardization. The solution was standardized periodically against standard 0.1 N potassium dichromate in the following manner. To 300 ml. of distilled water was added, with constant stirring, 2.5 ml. of sulfuric acid, 25.00 ml. of the standard dichromate solution and 2 grams of potassium iodide. The mixture was allowed to stand for 6 minutes in diffused light and then titrated with the thlosulfate solution, starch being used as the indicator.
0.01 N sodium thlosulfate solution. This reagent was prepared by diluting a measured amount of the aged and standardized 0.1 N sodium thlosulfate solution with freshly boiled and cooled distilled water. A few ml. of chloroform were added as a preservative* The solution was standardized daily in the following manner. To 300 ml. of distilled water was added, with constant stirring, 2.5 ml. of sulfuric acid, 25.00 ml. of the standard 0.01 N potassium dichromate solution and 0.5 gram of potassium iodide. The mixture was allowed to stand for 6 minutes in diffused light, and then titrated with the thlosulfate, starch being used as the indicator.


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Chlorine water. Chlorine gas was slowly bubbled, with occasional shaking, through 8 9 liters of distilled water contained in a black enameled bottle fitted with a siphon. Samples of the water were taken every few minutes, the chlorine concentration determined in accordance with the standardization procedure, and bubbling was discontinued when the chlorine concentration reached 1.1 1.2 milligrams per milliliter. When the atmosphere was excluded from the solution by a simple check valve the solution was found to retain Its strength for days, losing strength at a rather constant rate of approximately 0.015 milligrams per milliliter per day. Standardization was accomplished daily in the following manner. To 275 ml. of distilled water was added with swirling, 10 ml. of acetic acid and 0.75 gram of potassium iodide. 50.00 ml. of the chlorine water was measured beneath the surface of the solution, and it was titrated immediately with standardized 0.1 N sodium thlosulfate with starch as the Indicator.
Stock sulfide solution. 2 liters of distilled water were boiled at a moderate rate for 20 minutes in a 2 liter Erenmeyer flask and then cooled in a water bath under a nitrogen atmosphere. 7.1 grams of reagent grade NagS'OH^O crystals, freshly washed and blotted dry, were dissolved in the water, care being taken to prevent the entrance of air into the flask. The resulting solution had a sulfide concentration of approximately 0.5 milligram per milliliter.


- 18 -
and the concentration was found to remain fairly constant as long as the solution was stored under an atmosphere of nitrogen* The solution was standardised daily in accordance with the following procedure* 10*00 ml. of the sulfide solution were measured from a micro-buret beneath the surface of 25 ml. of sine acetate solution, and the resulting mixture was diluted with 110 ml* of distilled water* 60*00 ml* of 0*01 N iodine solution were pipetted in, the solution was acidified with 5 ml. of hydrochloric acid, and then it was allowed to stand for 6 minutes in diffused light* The mixture was titrated with standardised 0*01 N sodium thlosulfate, using starch as the indicator* A blank determination was carried out on the reagents*
Diluting water. The diluting water was prepared separately for each run by boiling 4.5 liters of distilled water for 45 minutes in a 5 liter, round-bottom, boiling flask and cooling in a water bath under an atmosphere of nitrogen* A water resulted that was practically free from oxygen, and it was stored under a pressure of nitrogen.
Standard buffer solutions. Clark and JLubs buffer solutions having pH values of 5.00, 7.00 and 0*60 were carefully prepared in accordance with the instructions given by Clark (49). These solutions were used in the investigation as standards by means of which the pH meter was calibrated at frequent intervals*
Concentrated buffer solutions* These buffers had to be


19 -
so constituted that the addition of a measured volume to a definite volume of diluting water would give a working buffer solution of certain known characteristics, and consequently their preparation presented many problems. Let us first mention the characteristics desired in the working buffer solutions, which may be defined as those solutions that were prepared in the reaction vessels prior to the addition Of the reactants to fix the conditions under which the reaction was to progress. It was thought that for convenience these solutions should have the same volume in every case throughout the investigation. The most convenient volume to work with was found to be 525 milliliters. Furthermore, It was considered that except in the case where the effect of ionic strength is the object of investigation the solutions should all have the same ionic strength where possible. Lastly, it was deemed necessary that the concentrations of buffer materials in the working solutions should be such that the pH value is never changed by more than about 0.05 pH unit upon the addition of the reactants.
It was found that the most convenient combination of volumes to use involved the addition of 25 ml. of concentrated buffer solution to 500 ml. of diluting water. Therefore, the concentrations of the concentrated buffers were based on this combination of volumes where ever possible. In the case of the pH 5 buffer solution the limited solubility of the buff er materials interfered with this plan, and a con-


- 20
centratea buffer of one-half the calculated strength was employed. This means that it was necessary to add 50 ml. of this particular concentrated buffer solution to 475 ml, of diluting water to give the previously decided upon 525 ml, of working buffer solution.
The materials used in the preparation of the concentrated buffers were those of the Clark and Xubs series (49), with potassium chloride being added in addition to equalise the Ionic strength values of the working buffers. Table 1 gives the compositions of the various concentrated buffer solutions employed during the course of the investigation as well as the ionic strength values and chloride concentrations of the working buffer solutions derived from these concentrated solutions. The solutions were made up in nitrogen-filled volumetric flasks with great care being taken to prevent the entrance of air. All water used In the solutions was rendered oxygen-free by boiling for some 30 -40 minutes and cooling under an atmosphere of nitrogen. The sodium hydroxide was added in the form of a carefully standardised, oxygen-free solution that had been stored under a pressure of nitrogen. The finished buffer solution was transferred to a nitrogen-filled storage bottle and was kept under a nitrogen atmosphere,
Some apprehension was felt over the possibility that the phthalate salts in the pH 5 buffer might be chlorinated during the runs involving this particular buffer. However,


TABLE 1
COMPOSITION OF CONCENTRATED BUFFER SOLUTIONS AND CHARACTERISTICS OF WORKING BUFFER SOLUTIONS.
pH of Components of concentrated buffer Ml/525 ml Characteristics of
solutions
working (grams/liter) working working buffer solutions
buffer KHC.H.O. KHJPO-, H_BO_ NaOH KC1 buffer *I without Added CI" I Witt
solution added KC1 added KC1
5.0 107.2 ~---- ........ 10w02 94.59 50 0.04885 2140 0,1093
6.0 7.184 26,87 25 0*09210 608 0.1093
6.4 ... 10.58 53*33 25 0*07520 1208 0.1093
6.6 14,95 37.05 25 0.08560 840 0,1093
6.8 19,70 19.36 20 0.09690 438 0,1093
7.0 ..... 12,44 tmrnrnmrnmam 25 0.05463 0.05462
7.2 ..... 114.4 .....: 23.52 20,77 25 0.09600 470 0.1093
7.4 26,54 9,492 25 0,1032 215 0,1093
8.0 ..; 58.96 .;.- 25 0.2154 :..... . 0.2154
9.0 ..... ..... 32.47 8,948 154.4 25 0.03565 3500 0.1093
* I refers to ionic strength value.


- 22
the cnlorine losses in the blank determinations with this buffer solution were of the same order of magnitude as those with the phosphate and borie acid solutions. Therefore, it is concluded that no error due to such a chlorinatien has been introduced.
Potassium chloride solutions. Potassium chloride solutions were employed in those parts of the investigation dealing with the effects of ionic strength and chloride concentration on the oxidation. They were used as a means of varying these two variables. For the ionic strength experiments a potassium chloride solution was desired of such strength that each 10 ml. substitution for diluting water in the 525 ml. volume of working buffer solution would Increase the ionic strength value of the working solution by 0.025 units. Such a solution was calculated to contain 97.86 grams of potassium chloride per liter.
In the case of the chloride experiments a potassium chloride solution was wanted of such concentration that 5 ml. substitutions for concentrated buffer solution (pH 7.0) in a series of working buffer solutions originally containing 50 ml. of concentrated buffer per 525 ml. of solution would yield a series of working buffer solutions having a constant ionic strength value of 0.1093 and chloride concentrations varying by equal increments from zero to approximately 2,000 parts per million. Such a solution was calculated to contain 85.52 grams of potassium chloride per liter.


-23
These potassium chloride solutions vere prepared in nitrogen-filled volumetrie flasks with a great deal of attention being devoted to the exclusion of .air. The distilled water used in the solutions was rendered oxygen-free by boiling for 30-40 minutes and cooling under nitrogen pressure. The finished solutions were transferred to nitrogen-filled storage bottles and were kept under an atmosphere of nitrogen.
Descrftptfofl of Apparatus
The reaction vessels used in this investigation were constructed from 625 ml., wide-mouth, amber glass, reagent bottles such as those commonly used in the packaging of laboratory chemicals. The lids are of the screw-cap variety and were molded from a plastic material. Conversion of the bottles to reaction vessels for the experiments involved the drilling of two holes, 6*5 mm. and 9.5 mm. in diameter, in each of the lids and the cutting of a gasket for each lid from a sheet of rubber packing about 2 mm. in thickness. The holes were drilled about 1 cm. from opposite edges of the lid end on a line passing through the center. They were fitted with rubber stoppers that had been carefully ground to the proper else to insure a positive closure. The gaskets were cut in the shape of a doughnut to fit snugly within the lids, but at the same time they allowed access to the interior of the vessels through the holes drilled in the lids. When the lids of the vessels were screwed down firmly against the gaskets


- 24 -
and the holes In the lids were stoppered the vessels were gas tight.
Figure 1 is a schematic diagram showing the nitrogen assembly that was used to fill the reaction vessels with nitrogen, to prevent the entrance of air into the vessels during the preparation of the working solutions and to maintain an atmosphere of nitrogen over the various stock solutions. The source of nitrogen was a steel cylinder (A) containing compressed nitrogen that was originally under a pressure of 2,200 pounds. The flow of gas from this cylinder was regulated by a needle valve (B) which was equipped with a pressure gauge. From the needle valve the nitrogen flowed through rubber connected glass tubing to a pressure regulator (C), to solution storage bottles (D) and to a manifold (E). The pressure regulator consisted of a 100 ml., ungraduated, glass cylinder that contained about two.inches of mercury and was stoppered by a 2-hole rubber stopper through which passed the nitrogen tube and the leg of a trap. The nitrogen tube extended to the bottom of the cylinder. This arrangement provided for a constant pressure of nitrogen in the system and allowed wasted nitrogen to escape without danger of air diffusing into the system. The storage bottles (D) were a series of 2-liter, wide-mouth bottles interconnected by rubber and glass tubing and containing concentrated buffer solutions, potassium chloride solutions and sodium hydroxide solution. They were: provided with rubber stoppers through which stoppered, glass


m 25
FIGURE I
SCHEMATIC DIAGRAM OF NITROGEN ASSEMBLY
\-2_kk
Legend A Sleet nitrogen cylinder B Nee die fa/re vr/th pressure gauge C Pressure regu/alor D So/ufion storage boll/e N i trogen man/fold F~ Reaction yessei
G Rack for supporting reaction resse/s H I Liter beaker


- 26
sleeves protruded. These sleeves were for the admittance of pipets so that the solutions could he withdrawn while a stream of nitrogen kept hack the air. The manifold (E) was constructed from 2.5 era. tubing about five feet in length, and it was fitted with nine, equally spaced, 0.5 cm. nipples to which rubber connections could be made. Two of these nipples served to connect nitrogen pressure to the diluting water and stock sulfide solution storage vessels. Connected to the remaining seven nipples were glass and rubber tubes that were used for filling the reaction vessels (F) with nitrogen. During the filling process the vessels were supported in an Inverted position on a wooden rack of simple design (G) which held seven of the jars. A 1 liter beaker (H) was used to catch the water forced from the reaction ressel by the nitrogen during the filling process.
Shewn in Figure 2 is a schematic diagram of the apparatus that was used in the preparation of the working solutions in the nitrogen filled reaction vessels. The diluting water was stored In a 5 liter, round bottom, boiling flask (A) which was provided with a 4 hole rubber stopper. Through the four holes in this stopper passed a nitrogen tube from the manifold (E), a rubber stoppered sleeve that served as a vent, a siphon to a volumetric plpet (B) and a return tube from the top of the plpet. This return tube closed the system and prevented exposure of the diluting water to air. The pipet (B) was a 500 ml., semi-automatic, volumetrio pipet especially designed


* 27 -
FIGURE 2
SCHEMATIC DIAGRAM OF APPARATUS USED IN
Leg end
A Diluting water storage tank 3 SOO ml. semi- automatic, yo/umelrlc pipet C Stock sulfide solution storage flask D 10 ml. micro buret ~ All troy en manifold F React ion vessel G -- Cnlor/hc water storage bottle
H SO ml. buret J 10 ml. micro buret


- 28
and constructed for this investigation, and it was calibrated to deliver 500 ml. at 25 degrees. Provision was made in the arrangement of the apparatus so that during the measuring of the diluting water into the reaction vessel (F) and during the subsequent additions of solutions a tube was available from the nitrogen manifold

29
bottle was connected to a 50 ml. buret (H) and a 10 ml. buret (X) by gum rubber connected glass tubing* The 50 ml. buret was of the semi-automatic type that is operated by a 2-way stopcock. The 10 ml. buret was a micro buret to which a side arm had been added just above the stopcock in order to convert it to a semi-automatic buret operated by a pinch-clamp.
The thermostat used in this investigation was of the conventional, manually operated type. The water bath consisted of a box 20 inches long, 13 inches wide and 11,5 inches deep that was constructed of 0,75 inch, cypress boards. The insulating properties of this wood are excellent, and the dimensions of the box allowed for a volume of water that was so large that the temperature of the water could be controlled with ease. Stirring in the bath was provided for by a turbine type stirrer located in one corner. Temperature control was accomplished by an arrangement whereby a portion of the water was circulated by means of a small circulatory pump through copper coils that were immersed in an ice bath. The proportion of water passing through the cooling coll was regulated with the aid of a by pass line. By careful adjustment of the system and constant vigilance the temperature variation could be held to plus or minus 0.1 C
The Beekman pH meter, model 0, was used for all pH measurements in the course of the Investigation. In the pH


30 -
range below 8,5 the normal glass electrode was employed, but at higher pH values the measurements were made using a high alkalinity, glass electrode.
Experimental Procedures
procedure was complicated by the necessity for keeping oxygen out of contact with the solutions and the various reagents as the solutions were being prepared. The steps were practically the same in all the experiments, but there were deviations that have been pointed out later in those cases where they have occurred. The first steps in the procedure were to fill the series of reaction vessels with distilled water and to screw on the lids in such a fashion that no bubbles remained in the jars. Glass tubes connected to the nitrogen manifold by short sections of gum rubber tubing and projecting through 1-hole rubber stoppers were then inserted into the larger of the two holes in each of the lids so that they extended to the bottoms of the jars. The smaller of the two holes in each of the lids was stoppered, and the vessels were inverted on a rack in the manner shown in Figure 1. The needle valve on the nitrogen cylinder was adjusted so that there was a rather vigorous bubbling of nitrogen through the mercury in the pressure regulator. In the case of each vessel in turn the stopper was removed from the smaller hole in the lid, the water was forced out and replaced by nitrogen, and the stopper


- 31
was replaced. After all the jars had been treated in this manner the stoppers were again removed, and a stream of nitrogen was allowed to flow through all of them simultaneously until no more than a drop or two of water remained in each vessel. The vessels were then restoppered, and the flow of nitrogen through the pressure regulator was reduced to a slow bubbling.
The next steps in the procedure deal with the preparation of the working buffer solutions in the nitrogen filled reaction vessels, and they were completed for each vessel once they were begun. A jar was removed from the rack, the nitrogen tube was removed, and a stopper was immediately put in its place. Now the vessel was placed in a position under the diluting water pipet (see B, Figure 1), the stopper was removed from the smaller of the two holes in the lid, and a nitrogen tube located near the pipet was immediately inserted to a depth of about one inch into the jar. After removal of the other stopper from the lid the leg of the pipet was inserted, and the jar was raised until the tip of the pipet leg was at the bottom of the vessel. 500 ml, of the oxygen-free diluting water were measured into the vessel from the plpet while a stream of nitrogen from the nitrogen tube prevented air from entering. Next the jar was lowered from the pipet and placed on the desk top, the stream of nitrogen still flowing from the nitrogen tube. 25 ml. of oxygen-free, concentrated, buffer solution were measured in by means of a pipet, and after re-


~ 32 -
moval of the nitrogen tube the vessel was restoppered. The jar was swirled several times to insure thorough mixing,, and then it was placed in the thermostat, where it was left for two hours to attain the desired temperature for the subsequent reaction.
After all of the reaction vessels of the series had been subjected to the treatment outlined in the above paragraphs they were ready for the introduction of the stock sulfide solutions and chlorine water. The addition of chlorine water followed very closely the addition of the stock sulfide solution in each vessel. However, in the interest of saving time a schedule governing these steps in the procedure was drawn up for each experiment, and in accordance with this schedule the reactions in some of the vessels of a series were completed and the results determined before the additions were even made to other vessels of the series. Consequently, the procedure for the additions is outlined as it was followed for an individual jar. The jar was removed from the thermostat, the smaller hole in the lid unstoppered and the nitrogen tube inserted. With a gentle stream of nitrogen flowing through the tube, the other stopper was removed from the lid, and the tip of the stock sulfide solution buret (see B, Figure 2) was inserted to a point beneath the surface of the buffer solution in the vessel. The volume of stock sulfide solution required to give the desired sulfide concentration, as calculated from the dally standardisation of the solution, was


33 -
measured into the vessel from the "buret. The buret tip was removed, and the solution was swirled several times to insure mixing* Now the tip of the 50 ml* or the 10 ml* chlorine water buret (see H or I, Figure 2), the sise of the buret depending upon the calculated volume to be added, was inserted beneath the surface of the solution, and the volume to be added, as calculated from the daily standardisation of the solution, was measured into the vessel as quickly as possible* At the same instant the chlorine water began to flow into the vessel the time was noted on a stopwatch so that the reaction time could be measured* The buret tip and the nitrogen tube were quickly removed, the vessel was re-stoppered, and thorough mixing was accomplished by swirling rapidly in small circles 12 15 times. The jar was replaced in the thermostat to await the previously decided upon time for stopping the reaction and determining its extent* The total time elapsed while the jar was out of the thermostat for the stock sulfide solution and chlorine water additions was 2 3 minutes*
Analytical methods. After the reaction time that had been previously decided upon for a particular vessel had elapsed, the extent of the reaction in that vessel was determined* The analytical method used in the determination depended upon whether the residual reactant in the vessel was chlorine or sulfide. The identity of the residual could usually be predicted from the original ratio of reactants and the conditions of the reaction. However, in the borderline


- 34
eases where it was difficult to predict what the residual would be the assumption was made that it was chlorine, and the procedure for the determination of residual chlorine was followed to the point where iodine either was or was not liberated from potassium iodide In the acid solution* Close observation at this point in the procedure showed whether the analysis was to be made for residual chlorine or residual sulfide*
The analytical method used for the quantitative determination of residual chlorine in the investigation was the standard iodometrio method for the determination of chlorine in water (50). The reaction in the vessel was stopped at the end of the reaction time by the addition of 0.75 grams of potassium iodide dissolved in enough acetic acid solution to lower the pH of the working solution to a value between 3 and 4. Usually 10 ml. of the acid was sufficient. After a thorough mixing by swirling, the solution was poured from the reaction vessel into an 800 ml. beaker, and the vessel was washed with a few portions of distilled water from a wash bottle* The liberated iodine was titrated with the standardized 0.01 N sodium thlosulfate solution, using 5 ml. of starch solution as the Indicator* The quantity of residual chlorine was calculated from the titration*
The method used for the quantitative determination of residual sulfide was based on the standard iodlmetrio method


35
for the determination of sulfides in water (51)* Enough acetic acid solution was added to lower the pH value of the solution to 3 or 4, and Immediately thereafter an excess of 0*01 K iodine was measured into the reaction vessel hy means of a pipet* These additions were made to the reaction vessel through one of the holes drilled in the lid, and every precaution was taken to prevent the loss of hydrogen sulfide from the jar* After thorough mixing the solution was poured into an 800 ml, beaker, and the reaction vessel was washed with several portions of distilled water from a wash bottle* The excess Iodine was titrated with the standardised 0*01 N sodium thlosulfate solution, using 5 ml* of starch solution as the Indicator* The amount of Iodine that reacted with the residual sulfide was calculated from the titration and a previously run blank that gave the relationship between the iodine solution and the sodium thlosulfate solution. The amount of residual sulfide was determined from the amount of iodine that reacts. *
Experiments to determine the effect of concentrations. The concentrations of sulfide selected for these experiments were those obtained by the addition of 1, 2 and 3 milligrams of sulfide to the working buffer solutions, these additions yielding concentrations of 1.9, 3.8 and 5.6 parts per million as sulfide or 2.0, 4.0 and 6.0 parts per million as hydrogen sulfide, respectively. Separate experiments were conducted at each sulfide concentration, and each experiment was made


- 36
with & series of seven or eight reaction vessels* Five of the vessels received the carefully measured amount of sulfide In accordance with the procedure previously described for th< preparation of -the working solutions, and the remaining vessels of the series were designated as blanks, receiving no sulfide* The chlorine dosages added to the five vessels containing the sulfide solutions in each experiment were 2, 4, 6, 8 and 10 times the sulfide dosage for the experiment* In the case of the vessels designated as blanks, chlorine dosages were added that were designed to give chlorine residuals of the same order of magnitude as those in the other jars of the series* Experiments dealing with each of the selected sulfide concentrations were made at pH values 5*0, 6.0, 7.0, 8.0 and 0.0. The reaction time and the temperature were held, constant throughout the experiments at 20 minutes and 25 C, respectively. It is pointed out that a deviation from the previously described procedure for the preparation of working solutions occurred in the case of the pH 5 solutions. In this particular case 25 ml. of the diluting water were withdrawn with a volumetric plpet from each of the reaction vessels, and 50 ml. of the concentrated buffer solution were added to the remaining 475 ml*, thus making a total volume of 525 ml. of working buffer solution as In the other cases.
Experiments to. de^rmftnft jthe. precision g a determination. In these experiments each reaction vessel in a series of four


- 37
vessels was treated in a manner as nearly identical as possible with the treatment received by all the other vessels of the series. Two milligrams of sulfide were accurately measured Into each of the vessels of a series in accordance with the procedure for the preparation of the working solutions* A chlorine dosage of 17,68 milligrams was added in each case, and the reaction time and the temperature were held constant at 20 minutes and 25 C, respectively, for all the experiments.. The experiments were run at pH values 5.0, 6.0, 7.0 and 8.0. The deviation previously mentioned in connection with the preparation of the working buffer solution at pH 5 also occurred during these experiments.
< ExTlsmiis to de.fremftne,,; jjte.^jffflfr ff#Mt- in-
action. These experiments were designed to give an indication of just how far the reaction progresses at the instant the reactants are brought into contact with each other. 0.75 gram of potassium iodide was dissolved in 5 ml. of starch indicator solution in each one of four 800 ml. beakers* Working buffer solutions were prepared in each of a series of four reaction vessels in accordance with the procedure previously outlined. These buffer solutions, each in its turn, were carefully poured into the 800 ml. beakers in such a manner that the dissolution of oxygen from the air was at a minimum, and four milligrams of sulfide were accurately measured beneath the surface of the solution. The tip of the 10-ml. chlorine water buret was inserted beneath the surface of the solution, and


- 38 -
the sulfide solution was Immediately- titrated with the chlorine water until a blue tinge appeared in the solution and persisted for a few seconds. This was taken as the end-point of the titration, the potassium iodide acting as an oxidation-reduction indicator in the presence of the starch. The experiment was conducted at pH values 5,0, 6.0, 7,0, 8.0 and 9,0, and the temperature was held constant at 25 C. for all runs. The deviation that has been mentioned concerning the preps-ration of the pH 5 working buffer solution occurred again during these experiments.
Experiments to determine the effect of time. The concentration of sulfide employed in these experiments was 3.8 parts per million as Sulfide or 4.0 parts per million expressed as hydrogen sulfide, which was the concentration that resulted when two milligrams of sulfide were added to the working buffer solution. Each experiment was made with a series of eight reaction vessels, two of which served as blanks containing no sulfide. Each of the remaining six vessels received two milligrams of sulfide in accordance with the procedure for preparing the working solutions, and then a dosage of 17.68 milligrams of chlorine was measured Into each one In Its turn. The ehlorlne dosage added to the blanks was one that was calculated to yield a chlorine residual comparable to those in the other six reaction vessels. The residual chlorine in each of the six vessels that were dosed with sulfide was determined at the end of 1, 5, 10, 20, 40 and 80


- 39
minutes, respectively. The same reaction times were employed in the case of the blank solutions. However, in order to get a complete blank run of six determinations covering the entire range of reaction times employed in the experimentsg it was necessary to combine blank determinations for three runs at the same pH value* Experiments were run at pH values 5*0, 6*0. 7.0, 8.0 and 9*0, while the temperature in all of the experiments was held constant at 25 C The deviation occurring in the preparation of the pH 5 working buffer solution Is again pointed out in connection with these experiments*
: /S&BEtaB&ft to. de^rmjne. Jhe. effect oX ..^mpjrjifrure,. Experiments identical to those described in the above paragraph were run at pH values 5.0* 6*0* 7.0, 8.0 and 9,0, but temperatures of 15 and 20 C. were substituted at each pH value for the ..250 C employed in the previous runs.
Experiments to determine the effect of ionic strength. These experiments were run in a manner similar to that used to determine the effect of time on the reaction. However, provision was made in the procedure for changing the ionic strength value of the working solutions from one experiment to another, and furthermore all the experiments were run at a constant pH value of 7,0, as nearly as the changing ionic strength would permit. A previously described potassium chloride solution was used to vary the ionic strength, The substitution of this potassium chloride solution for dilut-


- 40
ing water In the reaction vessels caused a deviation in the normal procedure for the preparation of the working solutions, the potassium chloride solution was substituted by 10 ml. increments; therefore, the diluting water was removed by 10 ml* increments with volumetric pipets in order to make room in the solution for the chloride solution so that the final volume of 525 ml* for the working buffer solution would remain unchanged, the ionic strength was varied from 0.05463 to 0.2046- during the experiments by substituting volumes of potassium chloride solution from 0 to 60 ml, in this fashion.
ffxitqr^nts; Jo. determine,, the, offecfr o chlorides, the chloride experiments were run in tne Same way that the Ionic strength experiments were run, with the exception that the deviations from the normal procedure for the preparation of the working solutions were employed to vary the chloride concentrations from one experiment to another, leaving the ionic strength value constant at 0.1093. All of the experiments were run at a pH value of 7*0 and a temperature of 25 C In order to vary the chloride concentration and leave the ionic strength value constant in the working buffers it was necessary to substitute potassium chloride solution of a stated strength for concentrated buffer solution in the preparation of the working buffer solutions* the potassium chloride solution used for this purpose has been described. In that experiment where no chloride ion was added, 50 ml. portions of concentrated buffer solution (pH 7,0) were added


- 41
to 475 ml. portions of diluting water obtained by withdrawing 25 ml. from each reaction vessel by means of a volumetric pipet. This resulted in a working buffer solution having an ionic strength value of 0.1093. In the other experiments this same Ionic strength value was maintained, but the chloride concentration was increased by regular increments by the substitution of 5 ml. increments of the potassium chloride solution for equal portions of the concentrated buffer solution in the working solutions. During the course of these experiments the chloride concentration was varied from 0 to about 2,000 parts per million by increasing the volume of the potassium chloride in the working buffer from 0 to 25 ml. and decreasing the concentrated buffer from 50 to 25 ml., the respective increases and decreases being made by 5 ml. increments.
pxperj,men,ts Ift.fffttftiffltflMS the, eftec^ o hydrMn-fton. concentration. No additional procedures were required in the study of the effect of hydrogen-ion concentration on the oxidation. The procedures used in the experiments concerned with concentrations and time were designed to lend information at the same time to the study of the effect of the hydrogen-ion concentration.


EXPERIMENTAL DATA
The results of the various experiments conducted during this investigation are expressed as ratios that show the number of units of ehlorlne reacted per unit of sulfide reacted. These ratios are given both in units of milligrams and moles* The milligram units are given because it Is easier for one to realise in milligrams than in moles the quantity of chlorine required to react with the sulfide, and it is a relatively simple process to calculate dosages for comparative purposes from data of this type. The ratios are presented in mole units for two reasons. In the first place, this manner of presentation shows more clearly the progress of the oxidation. It has been mentioned before that the only end-products of the reaction are free sulfur and sulfate produced by simultaneous reactions (46), Disregarding the actual mechanisms, the stoichiometric relationships involved in the production of these end-products may be represented as follows:
Cl2 / S 2C1- / S
4C12 / 8 / 4H20 s 8HC1 / S04
In accordance with these relationships the respective ratios
for the formation of free sulfur and sulfate are the followings
Moles of chlorine 1 Hole of sulfide
Moles of chlorine 4 Mole of sulfide
- 42 -


- 43
Since the products are formed simultaneously the ratio will generally assume some value between one and four* Therefore, considering that the ultimate end-product of the oxidation under the most favorable conditions must be sulfate alone, the magnitude of the value of the ratio may be taken as an indication of the progress of the reaction* The larger the ratio the greater is the quantity of sulfate that results from the oxidation*
The other reason for the use of the mole as the unit in expressing the ratio deals with a simple method for determln ing the amounts of free sulfur and sulfate formed by the reaction* Choppin and Faulkenberry (46) have indicated a method whereby the amounts of sulfur and sulfate produced can be calculated from a knowledge of the ratio in this form and the total quantity of sulfide oxidised* The method involves the use of two simultaneous equations developed from the above-mentioned stoichiometric relationships concerning chlorine and sulfide* Starting with one mole of sulfide and letting X mole go to sulfur and Y mole go to sulfate, one can obtain the equations,
X/Y: 1 (with respect to sulfide) and
X / 4Y s experimental ratio (with respect to chlorine),
which can be solved for X and Y once the experimental ratio has been established* The amount of free sulfur formed can be calculated in any desired units by multiplying X by the quantity of sulfide oxidised expressed in the same units*


- 44 -
Similarly, the amount of sulfate can he calculated by multiplying 2.99Y by the quantity of sulfide oxidised expressed in the desired units* Table 2 gives the values of X and Y calculated from certain values of the experimental ratio* Also included in the table are values for the ratio of moles of sulfur per mole of sulfate produced, which values illustrate the effectiveness of the experimental ratio as a means to indicate the progress of the reaction*
Precision of Determinations
The results of the experiments designed to indicate the precision to be expected in the overall treatment of each individual reaction vessel are presented in Table 3* It is seen from these results that the precision of the determination of the experimental ratios in the manner adopted for this investigation is dependent upon the hydrogen-ion concentration of the solution in whleh the oxidation takes place. Much greater precision is indicated in acid solutions than in solutions having pH values larger than 7.0. However, considering the numerous time consuming steps and measurements that must be taken in the preparation of the working solutions, the scant quantities of the reactants and the ever present possibility of some contamination by oxygen from the atmosphere, the precision of the determination appears to be satisfactory throughout the range of hydrogen-ion concentration included in the investigation.


TABLE 2
RELATIONSHIPS EXISTING BETWEEN CERTAIN VALUES FOR THE EXPERIMENTAL RATIO AND THE MOLES OF SULFUR AND SULFATE PRODUCED PER MOLE OP SULFIDE OXIDIZED.
Experimental Ratio r, Moles Sulfur per Mole of Sulfide (XV Moles Sulfate per Mole of Sulfide (Y) Moles Sulfur Mole Sulfate (x/r\
l.OO 1.00 0,000
1.10 0.967 0.033 29,3
1.20 0.934 0.066 14.2
1.30 0.900 0.100 9,00
1.40 0.867 0.133 6.52
1.50 0.833 0,167 5.00
1.60 0.800 0.200 4.00
1.70 0.767 0.233 3230
1.80 0.733 0.267 2.75
1.00 0.700 0.300 2.33
2.00 0.667 0.333 2.00
2.10 0.633 0.367 1.73
2.20 0.600 0.400 1.50
2.30 0.567 0.433 1.31
2.40 0.533 0.467 1.14
2.50 0.500 0.500 1.00
2.60 0.467 0.533 0.876
2.70 0.433 0.567 0,763
2.80 0.400 0.600 0,667
2.90 0.367 0.633 0.579
3.00 0.333 0.667 0,500
3.10 0.300 0.700 0.428
3.20 0,267 0.733 0,364
3.30 0.233 0.767 0.304
3.40 0,200 0.800 0,250
3.50 0,167 0.833 0.200
3.60 0,133 0.867 0,153
3.70 0,100 0.900 0.111
3.80 0,067 0.933 0.071
3.90 0,033 0.967 0.034
4.00 0,000 1.00 0.000
-45 -


TABLE S
PRECISION OP DETERMINATIONS AT VARIOUS pH VALUES
Volume of buffer solutions 525 ml.
Sulfide added} 2.00 mg.
Concentration of sulfide solutions 2 mg/529 ml or 3,78 ppm.
Chlorine addeds 17.68 mg.
Reaction times 20 min.
Temperature: 25 C,
pH Chlorine Reacted (ma) Chlorine to Sulfide Ratio (me) Chlorine to $ Sulfide Ratio (mole) Deviation from Average
4,03 15.60 7.80 3,53 0.3
15.55 7.78 3,52 0.0
15.58 7.79 3,52 0.0
15*57 ^ ** ... ^ ^
. 5.80 \ 12.20 6,10 2.76 0.4
:.- 12.09 6.05 2.74 -0.4
m 12.17 6.09 2.75 0,0
I2X08 w ...*.*.__ .H^R ^^B.
7.05 10.44 5.22 2.36 ; i,7'
10.10 5.05 2.28 -1.7
10.22 5.11 2.31 -0.4
10*26 _ 2..3J2..__mmm
7.92 8.35 4.18 1.89 0.0
* 8.51 4.26 1.93 2.1
m 8.35 4.18 1.89 0.0
m 8.24 4.12 1.86 -1.6
. 46


- 47 -
Effect of Concentrations on the Reaction
The experiments concerning the effect of the concentrations of the reactants on the oxidation were designed to Indicate the extent of the reaction under various conditions of concentration with the other conditions being held constant. It has been mentioned previously in the description of the procedure for these experiments that experiments were run at pH values varying by unit increments from 5,0 to 0,0 and that the reaction time and the temperature were held constant for all experiments. Attention is Invited to Table 1 for information relative to the ionic strength values and the chloride concentrations of the working buffer solutions In which the oxidations were carried out.
The results of the experiments are presented in Tables 4-6, inclusive, which deal with 1,90, 3,78 and 5,65 part per million sulfide solutions, respectively, It is seen from these tables that, within the limits of the experimental error involved in the procedure, the chlorine to sulfide ratio is Independent of the original concentration of sulfide in the range of sulfide concentrations investigated in the experiments. Furthermore, it appears that the results obtained with the 3,78 and 5,65 part per million sulfide solutions are more consistent than those with the 1.90 part per million solution. This is probably due in part to the greater accuracy that Is possible in measuring out the larger volumes of reagents re-


TABLE 4
EFFECT OF CONCENTRATIONS ON THE RATIO OF CHLORINE
TO SULFIDE REACTED. SULFIDE CONCENTRATION IS 1.90 PPM.
Volume of buffer solution: 525 ml.
Sulfide added: 1.00 rag.
Concentration of sulfide solutions 1 mg/527 ml or 1*90 ppm. Reaetlon times 20 min.
Temperatures 25 C.
pH Chlorine Added Imz) Chlorine Reacted (miE> Sulfide Reacted fmit) Chlorine to Sulfide Ratio Ima:} Chlorine to Sulfide Ratio (mole)
4.88 2.00 2.00 0.318 6.28 2.84
* 4.00 4.00 0.574 6.97 3.16
* 6.00 6.00 0.826 7.26 3.29
8.00 7.42 1.00 7.42 3.36
- 10.00 7.91 1.00 7.91 3.58
5.89 2.00 2.00 0.470 4.25 1.93
4.00 4.00 0.848 4.72 2.14
6.00 5.06 1.00 5.06 2.30
a 8.00 5.56 1.00 5.56 2.52
t 10.00 8.00 1.00 6.00 2.72
7.05 2.00 2.00 0.510 3.92 1.78
4.00 4.00 0.952 4.20 1.90
m 6.00 4.82 1.00 4.82 2.18
8.00 4.76 1.00 4.76 2.16
10.00 1.00 4.93 2.23
7.92 2.00 2.00 0.592 3.38 1.53
* 4.00 .... 1.00 .... ....
6.00 4.15 1.00 4.15 1.88
* 8.00 4.12 1.00 4.12 1.87
10.00 4.21 1.00 4.21 1.91
8.95 2.00 2.00 0.562 3.56 1.61
r 4.00 1.00 .... .....
m 6.00 1.00 .... '. ....
* 8.00 4.23 1.00 4.23 1.92
* 10.00 4.31 1.00 4.31 1.95
- 48 -


TABLE 5
EFFECT OF CONCENTRATIONS ON THE RATIO OF CHLORINE SULFIDE CONCENTRATION IS 3.78 PPM.
Volume of buffer solution: 625 ml*
Sulfide added] 2.00 mg.
Concentration of sulfide solution: 2 rag/529 ml er 3.78 ppm. Reaction times 20 min.
Temperatures 25 C*
pH Chlorine Chlorine Sulfide Chlorine to Chlorine to
Added Reacted Reacted Sulfide Sulfide
, (me) (me) Ratio (me) Ratio (mole)
4.90 4.00 4.00 0.694 5,77 2.62
H 8.00 8.00 1.18 6.78 3.07
12.00 12.00 1,70 7*06 3.20
16.00 14,08 2.00 7.49 3.39
* 20.00 16.041 2,00 8*02 3.63
5.89 4.00 4.00 0.901 4.44 2.01
m 8.00 8.00 1.61 4*97 2.25
n 12.00 11.04 2*00 5.52 2.50
m : 16.00 12.10 2.00 6.05 2.74
m 20.00 12,80 2.00 6.40 2.90
7.05 4.00 4.00 1.02 3.92 1,78
8,00 8.00 1.89 4,24 1.92
12.00 0.59 2.00 4.80 2.18
N 16.00 9.96 2.00 4.98 2.26
* 20.00 10.34 2.00 5.17 2.34
7.92 4.00 4.00 1.18 3.39 1.54
* 8.00 2.00
12.00 8.14 2.00 4.07 1.84
'fl- 16.00 8.15 2.00 4.08 1.85
it 20.00 8.28 2.00 4.14 1.87
8.95 4.00 4.00 1.13 3.54 1.60
* 8.00 2.00 ~
12.00 7.60 2.00 3.75 1.70
n 16.00 8.23s 2.00 4.12 1.87
n 20.00 8.27 2.00 4.14 1.87
- 49 -


TABLE 6
EFFECT OF CONCENTRATIONS ON THE RATIO OF CHLORINE
TO SULFIDE REACTED. SULFIDE CONCENTRATION IS 5,65 PPM.
Volume of buffer solution? 525 ml.
Sulfide added: 3.00 mg,
Concentration of sulfide solution! 3 mg/531 ml or 5,65 ppm. Reaction timet 20 min.
Temperature j 25 C.
pH Chlorine Chlorine Sulfide Chlorine to Chlorine to
Added Reacted Reacted Sulfide Sulfide
(mc) Cms) (mz) Ratio imp:) Ratio (mole)
4,90 6.00 6.00 1,01 5*94 2.69
* 12.00 12.00 1,74 6.90 3.12
18.00 18.00 2,51 7.17 3.25
24.00 22,82 3.00 7.61 3.45
30.00 24.12 3.00 8.04 3.64
5.89 6.00 6.00 1.27 4.72 2.14
12,00 12,00 2.29 5,24 2,37
18,00 16.01 3,00 5.64 2.55
# 24.00 17,72 3.00 5.91 2.68
30.00 18.30 3.00 6.10 2.76
7.05 6.00 6.00 1.50 4.00 1.81
12,00 12.00 2.76 4.34 1.97
n 18.00 14.15 3,00 4.72 2.14
e 24.00 14,85 3.00 4.95 2.24
30.00 15,61 3.00 5.20 2.36
7,02 6.00 6.00 1.74 3.45 1.56
i 12,00 **. 3.00 .... .....
it 18.00 12.06 3.00 4.02 1.82
24.00 12.26 3.00 4.09 1.85
** 30,00 12.65 3.00 4,22 1,01
8,95 6.00 6,00 1.66 3.62 1.64
12.00 -..... 3,00 .... ....
n 18,00 11.94 3.00 3.08 1,80
o 24.00 12,35 3.00 4.12 1.87
30.00 12.66 3.00 4.22 1.91
- 50


- 51
quired for the more concentrated solutions* At any rate the results for the 3*78 and 5.65 part per million solutions are summarised in Table 7, and the average ratios obtained from this summary are plotted in Figure 3. Several points are clearly illustrated by the family of curves in this figure. Probably the most striking feature Illustrated is the clear* cut dependence of the experimental ratio upon the hydrogen-ion concentration of the reaction medium. The ratios were quite large in a reaction medium having a pH value of 4.90, indicating that the larger portion of the oxidised sulfide was oxidised all the way to sulfate. However, the ratios decreased as the pH value of the reaction medium was raised, and it is apparent that limiting values were approached in alkaline media at a pH value somewhere in the vicinity of 8.0. The figure also brings into sharp focus the effect of the ratio of chlorine added to the sulfide on the experimental ratio and the dependence of the magnitude of this effect on the hydrogen-ion concentration of the reaction medium. The ratio of chlorine to sulfide reacted is seen to have increased in all eases with an increase in the ratio of chlorine added. This increase was quite substantial in acid solutions, but it became less marked as the pH value of the medium was increased. Finally, at a pH value in the neighborhood of 8.0 the characteristics of the increase in the experimental ratio with increased chlorine ratios appear to have become fixed and much less impressive than in acid solutions.


TABLE 7
EFFECT OF CONCENTRATIONS ON THE RATIO OF CHLORINE
TO SULFIDE REACTED. SUMMARY OF EXPERIMENTS
pH Mole Ratio of Chlorine Average ^Deviation
to Sulfide Reacted Ratio from
Sulfide Cone. Sulfide Cone* Average
3.78 ppm 5.85 ppra
(Tahle 5) CTahle 6)
4.90 2.62 2.69 2.66 1.3
3.07 3.12 3.10 0.8
n 3.20 3.25 3.23 0,8
m 3*30 3.45 3.42 0.9
m 3.63 3*64 3.64 0.2
5*89 2*01 2.14 2.08 3,1
* 2.25 2.37 2.31 2.6
a 2.50 2.55 2.53 1.0
m 2.74 2.68 2.71 1.1
* 2*90 2.76 2.83 2*5
7.05 1.78 1.81 1.80 0.9
1.92 1.97 1.95 1.3
2.18 2.14 2.16 1.0
* 2.26 2.24 2.25 0.5
2*34 2.36 2.35 0.5
7.92 a 1.54 1.56 1.55 0.7
1.84 1.82 1.83 0.6
1.85 1.85 1.85 0.0
1.87 1.91 1.89 1,1
8.95 1.00 1.64 1.62 1.3
MM. ..... .... ...
1.70 1.80 1.75 2.8
m 1.87 1.87 1.87 0.0
1.87 1.91 1.89 1.1
- 52 -


53 -
FIGURE 3
EFFECT OF CONCENTRATIONS ON THE RATIO OF CHLORINE
TO SULFIDE REACTED.
AVERAGE RAT/05 FOR 3.78 J 5.G5PPM. SULFIDE SOLUTIONS PLOTTED.
(See Table 7)


54
pH Chlorine Dosage
__im.....mxmM,MU$,}
4.90 6-8
5.89 4-6
7.05 4-6
7.02 2-4
8.05 2-4
Figure 3 capably illustrates another important consideration. It is noted from fables 4-6, inclusive, that In every case the quantities of chlorine that must be added to completely eliminate sulfide from the solutions are greater than the stoichiometric quantity required for oxidation to free sulfur, and these quantities are observed to depend to a great extent upon the hydrogen-ion concentration of the solution. Thus It is Indicated that there is an oxidation process whereby the oxidation of a portion of the sulfide Is carried beyond the free sulfur stage while there is yet unoxldlsed sulfide present in the solution. This is graphically shown in Figure 3, where It may be observed that the experimental ratio is always considerably greater than unity even when the ehlorlne is added on a mole to mole basis (2.2 milligrams per milligram of sulfide) or slightly less. From Tables 4, 5 and 6 it is determined that the chlorine dosages required to Insure the complete elimination of unoxidised sulfide in reaction media of various pH values lie between the following valuest


- 55
Effect of Time on the Reaction
It became evident from preliminary experiments concerning the rate of oxidation of sulfide solutions by chlorine that the Initial oxidation process is a very rapid one. This initial portion of the oxidation was found to give way to a much slower process* The experiments showed that the more rapid portion of the oxidation takes place within the first minute of the reaction, which period of time represents the smallest reaction time that can be conveniently allowed with the methods chosen for the investigation* It was further noted that the experimental ratio at the end of this reaction time was always considerably greater than unity, indicating that the oxidation at that point had proceeded beyond the sulfur state. Inasmuch as previous Investigators have suggested that sulfur is the primary product of the oxidation {46). it was decided that the study of the oxidation rate could best be made in two parts. The first part of the study was to deal with the Initial or rapid oxidation process that takes place within the first minute of the reaction. The procedure for an experiment designed to indicate the extent of the immediate reaction has already been described. In this experiment sulfide solutions of known concentration were titrated with standardised ehlorlne water under various known conditions, using potassium iodide in the presence of a little starch as an oxidation-reduction type Indicator,


56
This particular indicator was chosen because a study of the pertinent oxidation potentials reveals that it should indicate in acid solutions when the oxidation of the sulfide to sulfur has just been completed, and thus it should be possible to determine by its use whether or not it is probable that the immediate reaction Involves oxidation of the sulfide to free sulfur* the literature discloses that the above conclusion in regard to the usefulness Of potassium iodide as an indicator for this particular titration In acid solutions has been substantiated by experiment (52)* Further consideration of the problem indicated that the Iodide-iodine couple cannot be successfully employed in the titration of alkaline sulfide solutions with chlorine to indicate when the oxidation to sulfur is just completed* However, the experiments were conducted in the same manner in alkaline solutions with the thought that the results would still indicate the extent of the Immediate reaction under such circumstances* The results of the experiments ere presented in Table 8, It is evident from these results that the product of the initial oxidation process in acid solutions is free sulfur* Furthermore, it is just as evident that in alkaline solutions the immediate oxidation process carries the reaction beyond the free sulfur stage, the extent of the reaction depending upon the alkalinity of the medium within the limits chosen for this investigation*
The second part of the study of the effect of time on


TABLE 8
EXTENT OF THE IMMEDIATE REACTION BETWEEN CHLORINE AND SULFIDE.
Volume of buffer solution: 525ml.
Sulfide added: 4.00 mg.
Concentration of sulfide solution: 4 mg/533 ml or 7.51 ppm. Temperature: 25 C.
Reaction time: Sulfide solution is titrated with chlorine
water, using starch and El as the indicator.
pH Chlorine Chlorine to Chlorine to ^Deviation
Reacted Sulfide Sulfide from
(nut) Ratio (mz) Ratio (mole) Average
4.93 0.34 2.34 1.06 -0.5
* 9.31 2.33 1.06 -0.5
* 9.48 2.37 1.07 0.5
9.47 2.37 1*07 0.5
Average 1.065
5.89 8.86 2.22 1.01 0.0
* 8.87 2.22 1.01 0.0
8.89 2.22 1.01 0.0
8.89 2.22 1.01 0.0
Average 1.01 <
7.05 9*62 2.41 1.09 -0.9
9.77 2.44 1.10 0.0
9*83 2.46 1.11 0.9
* 9.84 2.46 111 . ' 0.9
Average 1.10
7.02 11.17 2.79 1.26 0.0
11.07 2.77 1.25 -0.8
11.13 2.78 1.26 0.0
11.23 2.81 lf27 0.8
Average 1.26
8.05 12.43 3.11 1.41 1.4
12.24 3.06 1.39 0.0
e 12.10 3,03 1.37 -1.4
Average 1.39


58
the oxidation was concerned with the reaction that takes place in the reaction mixture after the initial minute of reaction time* The procedure for the experiments that were invented to determine this effect has been described earlier* the sulfide concentration of 3,8 parts per million was selected for these experiments because it lies in a range of concentrations frequently encountered in natural waters, and it involved convenient quantities of reagents* Furthermore, it was demonstrated by the previously mentioned experiments concerning the effect of concentrations that under constant conditions the experimental ratio is independent of the original sulfide concentration as long as the latter is within the concentration limits considered in this investigation* The choice of 17*68 milligrams was adopted as the chlorine dosage to be added because it represents the stoichiometric quantity of chlorine that would oxidise the two milligrams of sulfide that were present in each of the solutions all the way to sulfate if the other conditions of the experiment would allow it* Furthermore, the earlier experiments dealing with concentrations showed that the chlorine dosage required to Insure complete removal of unoxidised sulfide in a reaction medium of pH value around 5*0 lies somewhere in the vicinity Of 8 milligrams of chlorine per milligram of sulfide* Attention is invited to Table 1 for information relative to the ionic strength values and the chloride concentrations of the working buffer solutions involved in these experiments*


- 59 -
The results of the experiments are presented in Table 9, It has been found from these results that when the common logarithm of the experimental ratio is plotted against the common logarithm of the reaction time a straight line is obtained. The evidence is presented in Figure 4, where the plots are shown for all of the experiments. The equations for these plots are of the type
log R Z log a / b log t,
or R at***
where R is the experimental ratio in mole units, t is the reaction time in minutes and a and b are constants. It Is seen that the constant a indicates the extent of the oxidation at the end of one minute of the reaction, and the constant b determines the slope of the line, or in other words it indicates the rate of the change in the experimental ratio with time. Thus the constant b may be regarded as an empirical rate constant. The constants derived from the results of each of the experiments are presented in Table 10, and it is observed that they are dependent to a very large degree upon the conditions under which the oxidation takes place* The experimental ratios calculated by the use of the constants are also exhibited in Table 10, and they are seen to be in good agreement with the observed ratios.
Effect of Temperature on the Reaction
The range of temperatures selected for use in the experi-


TABLE 9
EFFECT OF TIME ON THE RATIO OF CHLORINE TO SULFIDE REACTED.
OBSERVED DATA
PH Reaction Time (min) Chlorine Reacted Chlorine to Sulfide Ratio (hue) Chlorine to Sulfide Ratio (mole) Logarithm of Ratio (Mole)
4.90 1 14.92 7.46 3.38 0.529
m 5 15.05 7,53 3.41 0,533
10 15.27 7.64 3.46 0.539
20 15.30 7.85 3.47 0.540
* 40 15.46 7.73 3.50 0.544
80 15.44 7.72 3.50 0.544
5.89 1 11.38 5.69 2.58 0.412
' 5 11.92 5.96 2.70 0.431
n 10 11.96 5.98 2.71 0.433
n 20 12.20 6.10 2.76 0.441
. 40 12.38 6.19 2.80 0.447
m 80 12.50 6.25 2.83 0.452
7.05 I 9.72 4.86 2.20 0.342
5 9.60 4.80 2.18 0.339
10 10.24 5.12 2.32 0.366
20 10.46 5.23 2.37 0.375
.* 40 10.87 5.44 0.391
n 80 11.22 5.61 2.54 0.405
7.92 1 8.24 4.12 1.87 0.272
n 5 8.39 4.20 1.90 0.279
10 8.45 4.23 1.92 0.283
ft 20 8.38 4.19 1.90 0.279
40 8.75 4.38 1.99 0,299
9 80 8.70 4,35 1.97 0.295
8.95 1 7.60 3.80 1.72 0.236
5 7.99 4,00 1.81 0.258
10 8.09 4,05 1.83 0.263
20 8.26 4.13 1.87 0.272
40 8.41 4,21 1.91 0.281
80 8.53 4.27 1.93 0.286
- 00 -
Volume of buffer solutions 525 ml.
Sulfide added: 2.00 mg.
Concentration of sulfide solution: 2 mg/529 ml or 3.78 ppm.
Chlorine added: 17.68 rag.
Temperature: 25 C


- 61 -
FIGURE 4
EFFECT OF TIME ON THE RATIO OF CHLORINE TO SULFIDE REACTED
(See Table 9 )
055 t
i-1-1 i-1-> i-1-1-1-1
0 O.Z 0/t- D.& 0.8 1.0 I.Z 1.4 I.Q> 1.8 Z.O
Logorilhm of Time Min


TABLE 10
EFFECT OF TIME OK THE RATIO OF CHLORINE TO SULFIDE REACTED.
CALCULATED DATA
pH Reaction Constant Constant Calculated % Deviation
Time jt b Chlorine to from Observe*
(min) Sulfide Ratio
Ratio (mole)
4.90 1 3.39 0.00720 3.39 0,3
5 tt 3,43 0,6
10 tt 3.45 -0.3
B 20 m t 3.47 0,0
; 9 40 " 3.48 -0.6
m 80 m it 3.50 0.0
5.89 1 2.60 0.0193 2.60 08,:-:
5 2.68 -0,7
tt 10 2.72 0.4
20 t 2.76 0.0
40 ft 2.70 -0.4
80 it 2.83 0.0
7,05 1 2.04 0.0497 2.04 -7.5
m 5 tt tt 2.21 1.4
ft 10 tt tt 2.29 -1.3
20 tt 2.37 0,0
tt 40 tt 2.45 -0.4
80 ' tt 2.54 0.0
7,92 1 1,87 0.0124 1.86 -0.5
* 5 m 1.00 0.0
tt 10 ft w 1.02 0,0
tt 20 ?t 1.04 2.1
tt 40 t tt 1.95 -2.0
80 tt n 1,07 0.0
8.95 1 1.74 0.0243 1.74 1.2
5 t 1.80 -0.6
n 10 tt 1.84 0,5
tt 20 1.87 0.0
tt 40 1.00 -0,5
80 ft
- 62 -
Source of data for calculations: Table 9. Basis for calculations: R z at", where
R r Chlorine to sulfide ratio (mole),
Jt r Reaction time (min), : a and b are characteristic constants.


- 63 -
stents planned to Illustrate the effect of temperature on the oxidation extends from 15 to 25 centigrade. This range was chosen with the thought that it includes the temperatures of the great majority of sulflde-bearing waters. Furthermore, these temperatures were found to he the most convenient ones with which to work, using the equipment at hand. Temperature control with the thermostat constructed for use in this investigation was found to present a rather difficult problem when temperatures below 15 or In excess of 25 were employed. Also to be considered was the probability,, that temperatures much greater than 25 degrees were likely to yield doubtful results because of the adverse effect on the solubilities of the gaseous reactants. It has been mentioned that the procedure for the experiments used to determine the temperature effect was identical to that for determining the effect of time on the reaction. The time experiments were just repeated at 15 and 20 C., and the results were compared with the 25 degree results. *
The data observed during the 15 and 20 degree experiments are presented in Tables 11 and 12, respectively, and they are summarised in Table 13 along with the data from Table 9, which deals with Identical experiments conducted at 25 degrees. A study of the summary of the data reveals no pattern of change in the experimental ratios that can be attributed to the variations in temperature. Indeed, in the large majority of eases the maximum deviations between the ratios for the vari-


TABLE 11
EFFECT OF TEMPERATURE ON THE RATIO OF CHLORINE TO
SULFIDE REACTED. OBSERVED DATA
Volume of buffer solution: Sulfide added:
Concentration of sulfide solution:
Chlorine added:
Temperature:
525 ml. 2.00 mg. 2 rog/529 ml or 3.78 ppm.
17.68 mg. 15 C.
pH Reaetlon Chlorine Chlorine to Chlorine to
Time Reacted Sulfide Sulfide
{Min) Ratio (me) Ration (mole)
4.93 1 14.76 7.38 3.34
tt 5 15.17 7.59 3.44
10 15,24 7.62 3.45
tt 20 15.48 7.74 3.50
tt 40 15,56 7,78 3.52
m 80 15.68 7.84 3,55
5.97 1 11.36 5.68 2.57
6 12.05 6,03 2.73
* 10 12.02 6.01 2.72
tt 20 12.10 6.05 2.74
tt 40 12.38 6.19 2.80
* 80 12.43 6.21 2.82
7.14 1 9.41 4.71 2.14
5 9.93 4.97 2.20
tt 10 10.35 5.18 2.34
20 10.36 5.18 2.34
* 40 10.57 5,29 2.40
80 11.21 5,61 2.54
8.02 1 8.14 4.07 1.84
tt 5 8.23 4.12 1.87
10 8.52 4.26 1.93
' tt 20 8.42 4.21 1.91
40 8.64 4.32 1.96
tt 80 8.90 4.45 2.02
9.15 1 7.06 3.53 1.60
tt 5 7.76 3.88 1.76
tt 10 7.95 3.98 1.80
tt 20 7.99 4,00 1.81
tt 40 8.17 4.09 1.85
tt 80 8 .33 4,17 1.89
- 64 -


TABLE 12
EFFECT OF TEMPERATURE OK THE RATIO OF CHLORINE TO
SULFIDE REACTED. OBSERVED DATA
pH Reaction Time (rain) Chlorine Reacted (nuac) Chlorine to Sulfide Ratio (mx) Chlorine to Sulfide Ratio (mole)
4,96 1 14.96 7.48 3.39
n 5 15.05 7.53 3.41
n 10 15.27 7.64 3.46
20 15.40 7.70 3.49
40 15.54 7.77 3.52
80 15.33 7.67 3.48
5.92 1 11.35 5.68 2,57
5 12.02 6.01 2.72
10 11.98 5.99 2.71
m 20 12.12 6.06 2.75
40 12.34 6.17 2.80
m 80 12.54 6.27 2.84
7,09 1 9.83 4.92 2.23
* 5 9. 92 4.96 2.25
10 10.19 5.10 2.31
20 10,44 5.22 2.36
40 10.75 5.38 2.44
80 10.50 5.75 2,60
7,97 1 8.29 4.15 1.88
5 8.51 4.26 1,93
' 10 8*58 4.29 1.94
20 8.71 4.36 1.98
it 40 8.95 4.48 2,03
" 80 9.15 4.58 2.08
9.04 1 7.25 3.63 1.64
5 7.75 3.88 1.76
* 10 7.92 3.96 1.79
20 8.03 4.02 1.82
m 40 8.20 4.10 1.86
* 80 8.29 4.15 1.88
- 65 -
Volume of buffer solution! 525 ml.
Sulfide added: 2.00 mg.
Concentration of sulfide solutioni 2 rag/529 ml or 3.78 ppm.
Chlorine added: 17.68 mg.
Temperaturet 20 C.


TABLE 13
EFFECT OF TEMPERATURE ON THE RATIO OF CHLORINE TO
SULFIDE REACTED. SUMMARY OF OBSERVED DATA
Reaction : Time (mlnV Chlorine 15 to Sulfide (mole) 20 Ratios 25 ,Maximum Deviation
4.92 1 3.34 3. 39 3,38 1.5
H 5 3.44 3.41 3.41 0.9
10 3.45 3.46 3.46 0.3
0 20 3,50, 3.49 3.47 0.9
tl 40 3.52 3.52 3.50 0.6
0 80 3.55. 3.48 3.50 2.0
5.93 1 2.57. 2.57 2.58 0.4
n 5 2.73 2.72 2,70 1.1
0 10 2.72 2.71 2.71 0.4
ft 20 2.74 2.75 2,76 0.7
n 40 2.80 2.80 2.80 -0.0
# 80 2. 82 2.84 2. 83 0.7
7.10 1 2.14 2.23 2,20 4.1
it 5 2.20 2.25 2,18 3.2
t 10 2.34 2.31 2. 32 1.3
tt 20 2.34 2.36 2,37 1.3
40 2.40 2.44 2.46 2.5
ft 80 2.54 2.60 2,54 2.3
7.97 1 1.84, 1.88 1.87 2.2
tt 5 1.87 1.93 1.90 3.2
tt 10 1.93 1.94 1.92 1.0
20 1.91 1.98 1.90 4.1
40 1.96. 2.03 1.99 3.5
0 80 2.02 2.08 1.97 ^5.4
9.05 1 1.60 1.64 1.72 7.3
5 1.76 1.76 1.81 2.8
0 10 1.80 1.79 1.83 2.2
0 20 1.81 1.82 1.87 3.3
40 1.85 1.86 1,91 3.2
0 80 1.89 1.88 1.93 2.6
?The pH values listed here are the average values for the experiments conducted at the various temperatures*
66 -
Source of data: Tables 9, 11 and 12.


ous temperatures are probably well within the limits of experimental error for the methods employed in this investigation. Consequently it is believed that any temperature effect within the range of temperature used is so small that it cannot be determined by the investigative procedures applied, and it is therefore considered to be insignificant from the standpoint of this study.
Effect of Hydrogen-ion Concentration on the Reaction
It has been stated in the section treating of experimental methods that no additional procedures were required in connection with the study dealing with the effect of the hydrogen-Ion concentration on the oxidation, since those procedures used in the experiments concerning concentrations and time were designed to yield information at the same time about the hydrogen-ion effect, that the hydrogen-ion eoneen tration has a very decided effect upon the reaction is clear ly evident from an examination of the results of all the experiments presented thus far. However, in regard to this study only a portion of the previous data was considered.
The effect of the hydrogen-ion concentration on the extent of the Immediate reaction has already been noted, and for a review of the results of the experiments on this part of the investigation attention Is directed to Table 8. The remaining portion of the study concerns the hydrogen-ion effect in the reaction mixture after the reaction has been


68 -
allowed to progress for one minute and longer. In this regard particular consideration is invited to the data appearing in Table 9, the graphical representation of these data in Figure 4 and the related calculated data that are exhibited in Table 10. The essential facts may be quickly noted by referring to Figure 4. Sere it Is observed that the intercepts of the Log Ratio vs Log Time plots increase to a very pronounced degree with decreased pH value, which means that the extent of the reaction at the end of one minute varies considerably with the hydrogen-ion concentration* It is further observed that the slopes of the plots vary with the pH value in such a manner that there is a maximum slope indicated in the vicinity of pH 7,0. Since this slope is indicative of the reaction rate, it is seen that there is a possibility of correlating the reaction rate with the hydrogen-ion concentration. As a result of the above observations it was decided that the most satisfactory way to attack this phase of the study would be to demonstrate the effect of the hydrogen-ion concentration on the constants a and b of the empirical equation, R z atD, which has been found to be applicable to the reaction after one minute of reaction time. The significance of these constants has been pointed out in the discussion of the effect of time on the reaction.
Experiments such as those used to determine the effect of time on the reaction were found to be most suitable for


69
producing the necessary data for this part of the investigation. In fact the results for the time experiments conducted at pH values 4.90, 5.89, 7.92 and 8.95, which appear in Tables 9 and 10, were also employed in this study. The experimental conditions, other than pH value, in ell of these experiments were made as nearly Identical as possible* In addition to the above experiments similar experiments were conducted at pH values 6.35, 6.56, 6.79, 6.95, 7.25 and 7.44 in an attempt to determine the pH value at which the indicated maximal reaction rate occurs. The results of these additional experiments are offered in Tables 14 and 15, and all of the data in regard to the constants a and b that are used in the study are summarised in Table 16.
Figure 5 presents a very vivid, graphical Illustration of the effect of the hydrogen-ion concentration on the reaction that takes place in the period of time from one to eighty minutes. An examination of Curve I reveals that at a pH value of slightly less than 5,0 the reaction progresses during the first minute of reaction to a point that is quite near complete oxidation to the sulfate stage. The extent of this Initial reaction is then seen to decrease steadily as the pH value is increased until an apparent minimum is approached in the neighborhood of pH 9.0. The value of the experimental ratio at this minimum is evidently about 1.7, which means that the ratio of sulfur to sulfate produced Is about 3 moles to 1 according to the figures in Table 2.


TABLE 14
EFFECT OF HYDROGEN-ION CONCENTRATION ON THE RATIO OF CHLORINE TO SULFIDE REACTED. OBSERVED DATA.
Volume of buffer solution: 525 ml.
Sulfide addedi 2.00 mg.
Concentration of sulfide solution! 2 mg/529 ml or 3.78 ppm. Chlorine added: 17.68 mg.
Temperature: 25 C,
Ionic strength value: 0,1093
pH Reaction Chloride Chlorine Chlorine to Chlorine to
Time Concentration Reacted Sulfide Sulfide
(min) (nnmV (ma:) Ratio (mk) Ratio (mole)
6.35 1 1200 10,66 5.33 2.42
it 5 10.99 5.50 2.49
10 m 11.20 5.60 2,54
20 9 11.45 5.73 2.60
* 40 11.96 5.98 2.71
80 ft 12.14 6.07 2.75
6.56 1 832 10.07 5.04 2.28
ft. 5 10.81 5.41 2.45
it 10 ft 10.84 5.42 2.46
n 20 11.57 5.78 2.62
ft 40 tt 11.96 5.98 2.71
* 80 -w 12,33 6.17 2.80
6.79 1 135 9.96 4.98 2,26
* 5 10,71 5.36 2.43
10 10.80 5.40 2.45
20 11.46 5.73 2,60
40 11.84 IS #05* 2.68
80 12.18 6.09 2.76
6*95 1 1540 9.52 4,76 2.16
n 5 10,64 5.32 2,41
ft 10 tt 10.39 5,20 2,36
20 it 10.98 5.49 2,49
ft 40 ft 11.60 5.80 2,63
ft 80 11.88 5,94 2,69
Continued
- 70 -


TABLE 14
Continued
pH Reaction Chloride Chlorine Chlorine to Chlorine to Time Concentration Beaeted Sulfide Sulfide
^min) 7.25 1 466 8.93 4.57 = 2.07:.
n 5 9.60 4.80 2.17<
10 B 9.68 4.84 '2.19
20 #t 10.06 5.03 2.28
t 40 10.52 5.26 2.38
ft 80 ft 10.92 5.46 2.48
7.44 1 213 9.00 4.50 2.04
* 5 -.tt 9.33 4.67 2.12
. 10 9.20 4.60 2.08
i 20 it 9.96 4.98 2.26
ft 40 ft 10.18 5.09 2.30
# SO 10.39 5.20 2.36
71


TABLE 15
EFFECT OP HYDROGEN-ION CONCENTRATION ON THE RATIO OF CHLORINE TO SULFIDE REACTED* CALCULATED DATA.
Source of data for calculationss Table 14, lasls for calculations: R z at", where R s Chlorine to sulfide ratio (mole), Jt r Reaction time (min), a and b, are characteristic constants,
pH Reaction Constant Constant Calculated % Deviation
Time
Chlorine to from Observed
(rain) Sulfide Ratio (mole) Ratio
6.35 1 2.38 0.0333 2.38 -1.7
5 2.51 1.2
10 ft 2.57 1.2
* 20 ft tt 2.63 1.1
* 40 2.69 -1,1
* 80 ft 2,75 0.0
6.56 1 2.25 0*0502 2.25 -1.3
n 5 ft 2.44 -0.4
ft 10 * it 2.52 2.4
tt 2G 2.61 -0.4
ft 40 tt 2.70 -0.4
ft 80 it ft 2,80 0.0
6.79 1 2,24 0.0473 2.24 -0.9
5 it m 2.42 -0.4
10 2.50 1.0
20 2.59 -0.4
40 ft 2.67 0,4
80 2.76 0,0
6.95 1 2.16 0,0502 2.16 0,0
5 2.34 2'. 9
ft 10 ft' ft 2.43 2.9
20 tt 2.51 0.8
40 ft 2.60 -1.1
it 80 2.69 0.0
Continued
72 -


TABLE 15
Continued
pH Reaction Constant Constant Calculated % Deviation
Tine a b Chlorine to front Observed
(min) Sulfide Ratio
Ratio Craole)
7.25 1 1.99 0.0499 1.99 -3,9
* 5 tt 2.16 -0.5
10 tt 2.24 2.3
20 2.31 1.3
tt 40 t tt 2.46 0.8
80 tt tt 2,48 0.0
7.44 1 2.03 0.0349 2.03 -0.5
tt 5 tt 2.14 0.9
10 ft tt 2.19 5.1
* 20 tt tt 2.25 -0.4
tt 40 it 2.30 0.0
tt 80 tt 2.30 0.0
- 73 -


TABLE 16
EFFECT OF HYDROGEN-ION CONCENTRATION ON THE CONSTANTS OF THE EMPIRICAL EQUATION, R = at0. SUMMARY OF DATA,
pH Ionic Strength Value Chloride Concentration (nam) Constant a Constant b
4.90 0.1093 2120 3.39 0.00720
5.89 0.1093 603 2.60 0.0193
6.35 0.1093 1200 2.38 0.0333
6.56 0.1093 832 2.25 0.0502
6.79 0.1093 435 2.24 0.0473
6.95 0.1093 1540 2.16 0.0502
7.25 0.1093 466 1.99 0.0499
7.44 0.1093 213 2.03 0.0349
7.92 0.2154 0 1.87 0.0124
8.95 0,1093 3470 1.74 0.0243
- 74 -
Source of date: Tables 1, 10, 14 and 15,
Concentration of sulfide solutions: 2 mg/529 ml or 3.78 ppm. Chlorine added: 17.68 mg.
Reaction time: 1-80 minutes.
Temperature: 25 C.


- 75 -
FIGURE. S
EFFECT OF HYDROGEN- ION CONCENTRATION ON THE CONSTANTS OF THE EMPERICAL EQUATION', R = atb.
CSee Table IG>)
3.8 -
Car ye I: Constant a-Curve U : Constant t>
-0.054
D.OOG


- 76 -
Curve IX of Figure 5 demonstrates that there Is a maximum oxidation rate for the reaction mixture that is very sharply defined and dependent upon the hydrogen-ion concentration* This maximum Is very broad, extending between the approximate pH values 6.5 to 7.3, and the rate constant is seen to drop off quite steeply on both sides* Toward the acid side the constant attains a considerably lower value than on the alkaline side, and it appears to be destined for a value of zero at a pH somewhere around 3.5. This is to be expected since Curve I indicates that the oxidation goes completely to sulfate within the first minute of reaction time at a pH value in the general vicinity of 4.0. On the alkaline side of the maximum the constant levels off sharply at pH 9.0, approximately, while it still has a value that is about one-half that at the maximum. Attention is called to the fact that the value of constant b for the experiment conducted at pH 7.92 has been disregarded in drawing Curve II. Table 16 discloses that the solutions used in this particular experiment were calculated to have ionic strength values of 0.2154 as compared to values of 0*1093 for the solutions employed in the other experiments of this study* This difference In lonle strength values could not be avoided In the investigative methods used, and later work concerning the effect of ionic strength on the reaction justifies the disregarding of the above-mentioned value.
In Figure 6 some of the experimental data from Tables 9


,* 77 -
FIGURE 6
EFFECT OF HYDROGEN-ION CONCENTRATION ON THE
RAT/0 OF CHLORINE TO SULFIDE REACTED
( See Tables 9 f 14 )
Curve I- Reach.au time is 1 mirt.
Curve U: Reaction time is 10 min.
Curve UJ' Reaction time is 40 min.
3.6
\, 34


c 32
3.0

^ 2.8
^ 2.6
* 24
^ 2,2 Xn \ V -

^ 2.0 -O m
1.8 -on
1.6 oi 1
5.0 G.O 7.0 8.0 ptt 3.0


- 78 -
and 14 are presented In a way designed to illustrate with greater clarity some of the existing relationships that are not readily apparent from an examination of Figure 5, The experimental ratios obtained with reaction times of 1, 10 and 40 minutes, respectively* have been plotted against pH values to give an Interesting family of curves. Curve I of Figure 6 and Curve I of Figure 5 may be regarded as being the same, each of them showing how the experimental ratio for a one minute reaction time varies with pH, However, in the case of Figure 5 the curve was derived from calculated data, whereas in Figure 6 all of the curves represent the observed data. Thus what has been said about Curve I of Figure 5 may be applied to Curve I of Figure 6* In fact* with a few notable exceptions the same general description is applicable to Curves II and III, which represent the experimental ratios obtained with reaction times of 10 and 40 minutes, respectively. It is observed that these latter curves are displaced upward in relation to Curve I, which is to be expected when the longer reaction times are considered. The most striking feature of the two curves, however, is the bumped area in each of them in the vicinity of pH 7,0, The depth of the bumped area is seen to be a function of the reaction time, which further illustrates the maximal oxidation rate in this pH range, A significant feature of this maximum in the oxidation rate is readily recognised from the relationships that exist between the curves presented


- 79
In Figure 6. It is apparent that in the region of the bumped area the oxidation may proceed at a higher pH value with a minimum of additional time to the same point that is indicated by a considerably lower pH value and a one minute reaction time* For example, an experimental ratio of about 2.8 is indicated for a one minute reaction time at a pH of about 5,9, but this same ratio can be attained at a pH value of about 7,0 by allowing a reaction time of 40 minutes. Were it not for the maximal oxidation rate in this pH range it would require a much greater reaction time to attain the given ratio at the pH value of 7,0,
The manner in which the experimental ratio increases with the hydrogen-ion concentration, as illustrated in Figure 6, is suggestive of the way in which the concentration of undlssociated hypoohlorous acid must increase with hydrogen-ion concentration in a solution of chlorine water. If an analogy between these two relations could be drawn it might give some indication as to whether or not free hypoohlorous acid is the effective oxidizing agent in the reaction. Consequently* calculations have been made according to the method suggested by McKinney (53) to Illustrate the relationships that exist between the activity fractions of free hypoohlorous acid and the hypochlorite ion and the hydrogen-ion concentration. The method that is Involved is as follows,


- 80 -
For hypoohlorous acid -
HC10 fi^ / CIO"* K Z 5.6 x 10~8 (29)
aHC10
*** aHCI0 5 1
2
aH,
s, t a, si : 5.6 x 10"8 HC10 CIO BT ^ J
a^ 2 10-PH
a
fiCIO CIO
'Wo 1 acio- 2 1 1 * 10(58 8>
Ijet a a 2 a
HC10 r CIO T
aBCl(/a = 1 / 5.6 xl<&* ~8> 1 / 10PH 7.252)
a,- 5.6 10 liyCgH 7*252)
C1 ^ 1 /s.6 k> -8) 1 / 10^ 7-252) The calculated activity fractions are presented In Table 17. and the data are exhibited graphically in Figure 7. A comparison of the hypoohlorous acid curve. I, of Figure 7 with the ratio curves of Figure 6 reveals some relationships that appear to be rather significant. It is very Interesting to note in this connection that in the pH range from 9.0 to 10,0, where free hypoohlorous acid is observed to be practically non-existent, the experimental ratio for a given reaction time is apparently at a minimum. As the pH decreases from


TABLE 1?
ACTIVITY FRACTIONS HYFOCHLOROBS ACID
pH aHC10/aT 8C10VeT
3.2 1.0000 0.0000
3*4 0.9999 0.0001
3*6 0.9998 0.0002
3.8 0*9996 0.0004
4.0 0.9994 0.0006
4.2 0.9991 0.0009
0.9986 0,0014
4.6 0.9978 0.0022
4.8 0.9965 0.0035
5.0 0.9944 0.0056
5,2 0.9012 0.0088
5.4 0.9861 0.0139
5,6 0.9782 0.0218
5.8 0.9659 0.9470 0.0341
8,0 0.0530
6.2 0.9185 0.0815
6,4 0.8764 0,1236
6.6 0.8177 0.1823
6.8 0.7391 0.2609
7.0 0.6410 0.3590
7.2 0.5299 0.4701
7.4 0.4149 0.5851
7.6 0.3096 0.6004
7.8 0.2208 0.7702
8.0 0,1515 0.1013 0.8485
8.2 0.8987
8.4 0.0662 0.9338
8.6 0.0429 0.9571
8.8 0.0276 0.9724
9.0 0.0175 0.9825
9.2 0,0112 0.9888
9.4 0.0070 0.9930
9.6 0.0045 0.9955
9.8 0,0028 0,9972
10.0 0.0018 0.9982
10.2 0.0011 0.9989
10,4 0.0007 0.9993
10,6 0.0005 0.9995
10.8 0,0003 0.9997
11.0 0.0002 0.9008
11.2 0.0001 0*9999
o.opoo 1-0000
81 -


82 -
FIGURE 7
ACTIVITY FRACTIONS- HYPOOHLOROUS ACID
(See Table 17)
Curve!: Activity fraction of undissociated hypochlorous acid. Curve TL: Activity fraction of hypochlorite ion.


83 -
this vicinity the activity fraction of toe hypoehlorous acid and the experimental ratio both increase sharply, and the curves are quite similar in appearance in the alkaline region, the similarity being particularly striking in the case of Curve III of Figure 6* However, In the more acid range the curves bear little resemblance to one another* The activity fraction curve of hypoehlorous acid begins to taper off at pH 6.5 and to approach a maximum value of unity; whereas the experimental ratios at this pH value are increasing even more sharply with decreased pH.
Effect of Ionic Strength on the Reaction
The previous observations concerning the effect of hydrogen-ion concentration on the reaction serve to emphasise toe necessity for rather severe pH control throughout the various experiments included In the investigation of this oxidation* It has been mentioned in connection with the preparation of the concentrated buffer solutions that the maximum allowable change in pH value of any working solution due to the addition of reagents and any subsequent reaction was about 0*05 pH unit, and it was found that considerable concentrations of buffer materials were required in the working solutions to limit the pH change to such narrow limits* Consequently, it was necessary to study the oxidation in buffer solutions having relatively large ionic strength values* It has been seen in the preceding work that an ionic


- 84 -
strength value of 0.1093 was adopted as a convenient and constant value to he used in most of those experiments where the use of solutions having such an Ionic strength was possible. However, Information was desired relative to the reaction in solutions having ionic strength values comparable to those values found in natural waters, such values being in the general vicinity of 0.005. In order that such information could be obtained, experiments designed to Indicate the effect of ionic strength on the reaction were conducted over a wide range of ionic strength values with the thought that the results could be extrapolated to lesser ionic strength values, the procedure for these experiments has been described earlier.
It has been pointed out In the description of the procedure for tee experiments dealing with the effect of Ionic strength on the reaction that the Investigative methods used were similar to those used in determining the effect of time. The notable difference was that a means was provided to vary the ionic strengths of the working solutions between the individual experiments while leaving the pH value constant. All of the experiments were conducted in solutions having a pH value that was as nearly constant as possible at 7.0, This value was chosen because It lies at the middle of the range of pH values considered in this investigation, and it is possible to attain a lower Ionic strength value in solutions at this pH value than others with the assurance


85 "
that the pH change due to the addition of reagents will he kept within the prescribed minimum of 0,05 unit* Furthermore ,,; it can he seen from Figure 5 that the selected value also lies within a range of pH values where the effect due to small variations in the hydrogen-ion concentration is practically negligible*
The results observed during the ionic strength experiments are shown in Table 18* It is noted that the pH value decreased steadily from 7*05 to 6*89 as the ionic strength was Increased from 0*05463 to 0.2046 during the experiments-However, the previously accomplished work concerning the effect cf hydrogen-ion concentration indicates that this small variation of 0.16 pH unit produces no inaccuracies due to that effect in determining the effect of ionic strength on the reaction as long as the experiments are conducted within the range of pH values lying between 6.5 to 7.3.
The constants, a. and b, of the empirical equation, R = at*, were calculated for each of the experimental runs in the study, and the results are reported in Table 19. An examination of toe results indicates that the Initial In-crease in ionic strength apparently produced a marked increase in the value of constant a, but further increases in the ionic strength value had no recognizable effect on this constant. A glance at Table 18 shows that the initial increase in Ionic strength was accompanied by a change in the chloride concentration of the working solutions from 0 to


TABLE 18
EFFECT OF IONIC STRENGTH ON THE RATIO OF CHLORINE TO SULFIDE REACTED. OBSERVED DATA.
Volume of buffer solution. 525 ml.
Sulfide added! 2.00 rag.
Concentration of sulfide solution! 2 mg/529 ml or 3.78 ppm Chlorine added! 17.68 mg.
Temperature t 25 C.
Ionic Reaction pH Chloride Chlorine Chlorine to Chlorine to
Strength Time Concentration Reacted Sulfide Sulfide
(min) (mnn) Ratio (mis) Ratio (mole)
0.05463 1 7.05 0 9.72 4.86 2.20
* 5 41 tt 9.60 4.80 2.18
10 ft 10.24 5.12 2.32
20 ft ft 10.46 5.23 2.37
40 ft m 10.87 5.44 2.46
tt 80 ft 11.22 5.61 2.54
0.07963 1 7.02 880 10.21 5.11 2.32
w 5 *. 10.57 5.29 2,40
ft 10 ft 11.50 5.75 2.60
ft 20 ft tt 11.85 5.93 2.68
ft 40 ft ft 12.31 6.16 2.79
80 ft ft 12.79 6.40 2.90
0.1046 1 6.98 1760 10.23 5,12 2.32
5 tt 10.42 5.21 2.36
* 10 ft tt 11.09 5.55 2.51
ft 20 ft 11.20 5.60 2.54
40 11.82 5.91 2.68
* 80 ft ft 12*22 6.11 -: 2.77 ,..


TABLE 18
Continued
Ionic Reaction pH Chloride Chlorine Chlorine to Chlorine to
Strength Time Concentration Reacted Sulfide Sulfide
(rain) (nam) {ma* Ratio (nut) Ratio (mole)
0.1293 1 6.95 2640 10,09 5.65 2.29
* 5 it 10.20 6.10 2.31
10 it 10,95 5.48 2.48
20 11.13 5.57 2.52
40 it 11.32 5,66 2.56
* 80 tt .......*........ 11.64 5.82 2.64
0,1546 1 6.93 3520 9.96 4,98 2.26
tt 5 , tt 10.47 5.24 2.37
tt 10 tt t 10.83 5.42 2.46
20 tt tt 11.13 5.57 2.52
40 it tt 11.15 5.58 2.53
R 80 tt 11.48 6,74 2.60
0.1796 1 6,91 4400 10.09 5.05 2.29
5 , 10.34 5.17 2.34
10 ; 10.72 5.36 2.43
20 10.90 6.45 2.47
40 tt 10.94 5.47 2.48
80 tt 11,13 5.57 2.52
0.2046 1 6.89 5280 10.20 5.10 2.31
it 5 it 10.16 5.08 2.30
10 tt: 10,73 5.37 2.43
tt 20 tt ft 10.85 5.43 2.46
: 40 , . tt 11,09 5.55 2.51
tt 80 n. 11.29 5.65 ,,,,,-,' 2,56...............'.....


TABLE 19
EFFECT OP IONIC STRENGTH ON THE RATIO OF CHLORINE TO SULFIDE REACTED. CALCULATED DATA.
Source of data for calculationst Table 18. Basis for calculations: R 1 at", where
R z Chlorine to sulfide ratio (mole),
t"a Reaction time (min),
a and b are characteristic constants.
Ionic Strength
Reaction Time (rain)
Constant a
Constant 3l
Calculated Chlorine to
Sulfide Ratio (mole)
% Deviation from Observed Ratio
0*05463 HI 2.04 0*0497 2.04 -7.5
5 it 2.21 1.4
10 it 2.29 -1.3
20 w ii 2.37 0.0
40 ii 2.45 -0,4
80 ii 2.54 0.0
0.07963 1 2*30 0.0526 2.30 -0,9.
5 ii 2,51 4.5
10 if 2.60 0.0
20 2.70 0.7
40 # 2.80 0.4
it 80 tt 2*90 0.0


TABLE 10
Continued
Ionic Strength Reaction Time (min) Constant a Constant H Calculated Chlorine to Sulfide Ratio (mole) % Deviation from Observed Ratio
0*1046 1 2.23 0.0492 2.23 -3.9
w 5 tt 2.42 2.5
it 10 tt 2.50 -0.4
tt 20 tt 2.59 1.9
tt 40 tt ft 2.68 0.0
* 80 it. tt 2.77 0.0
0.1296 1 2.2S 0.0331 2.28 *0.4
* 5 it tt 2.41 4.2
it 10 tt 2.47 -0.4
tt 20 tt tt 2.52 0.0
it 40 tt tt 2.58 0.8
tt 80 tt tt' 2.64 0.0
0.1546 1 2.27 0.0307 2. 27 0.4
it 5 2.39 0.8
tt 10 tt 2.44 -0,8
it 20 it tt 2.40 -1.2
tt 40 it tt 2.55 0.8
80 tt 2.60 0.0


TABLE Id
Continued
Ionic Reaction Constant Constant Calculated % Deviation
Strength Time a Chlorine to from Observed
(rain) Sulfide Ratio
Ratio (mole)
0,1796 1 2, 28 0.0238 '2 28 -0,4
tt 5 2,37 1.3
' ft 10 "tt ** 2.41 -0,8
ft 20 ft 2,45 -0.8
ft 40 tt 2.40 0,4
. 80 2.53 0.4
0,2048 HI 2,29 0,0251 2,29 -0,9
* 5 it n 2,39- 3.8
10 ft 2,43 0.0
20 n 2.47 0.4
40 n 2.52 0,4
* 80 # 2,56 0.0


91 -
880 parts per million. Literature previously cited (42) suggests the possibility that the observed effect on constant a. may be due to this addition of chloride, and this point was investigated further by later experiments planned to demonstrate the effect of chlorides on the reaction* Constant b was observed to exhibit a definite variation vlth changing Ionic strength. The relationships that exist between the two are represented graphically in Figure 8. It appears that b has a constant value in the vicinity of 0.050, within the limits of experimental error, for ionic strength values 0.05463, 0,07963, 0.1046 and the previously studied value of 0.1093. However, a sharp decrease occurs in the value of the constant with Increasing ionic strength values just in excess of this range, until it appears that a minimum value of about 0.022 to 0,023 is approached at an Ionic strength value somewhat greater than 0.20. An extrapolation of the data shown In Figure 8 to the range of ionic strengths normally encountered in natural waters suggests that the reaction is unaffected by variations in ionic strength in this region. In fact, it is indicated that as far as the effect of ionic strength is concerned, the results of most of the preceding experiments are applicable to such dilute solutions as those found in natural waters, It is pleasing to note that this study explained the apparent discrepancy In the value of the constant b for pH value 7.92 in Curve II of Figure 5. All of the other points on this curve were de-


- 92 -
FIGURE 3
EFFECT OF ID n:c strength on the constants of
THE EMPIRICAL EQUATION, R r*. {See Table 19)
0.070
0.065
0.060
0055
0.050 O

00+5
0.040
0.015 O \
0.030 \ o
0.025 0.020 L i i 1 II 1 1 1 1 1 1
1 o.oz 0.04 0.06 0.08 \0.I0 CIS. 0.14 6.1 & 0.18 O.Z0 OZZ Ionic Strength


- 93 *
termined In solutions having a common ionic strength value of 0,1093, but this particular point in question was determined at a value of 0,2154, It is seen from Figure 8 that decreasing the ionic strength from 0,21 to 0,11 has the effect of increasing the value of constant b by a factor of 2,2, Applying this factor to the point in question in Figure 5, it is seen that the value obtained for the constant is brought into very good agreement with the presupposed value.
In view of the foregoing evidence that the effective agent in this oxidation reaction is hypoehlorous acid, it was thought advisable to investigate the effect of the concentration of chloride ion, which was used to vary the ionic strength values during these experiments, on the hydrolysis of chlorine to hypoehlorous acid. The equation for this hydrolysis reaction is as followst
CI" / HC10 / Cl2 / H20 Latimer (29) gives -6,315 calories per mole as the free energy change for this reaction, which yields an equilibrium constant of 4,27 x 104. Using this value for the equilibrium constant, it is calculated that at a pH value of 7,0 and with no chloride added all of the chlorine that was added to the working solutions involved in the experiments existed in the solutions as hypoehlorous acid. This acid was Ionised in accordance with the relationships shown in Figure 7, Further calculations show that when toe chloride concentration in the


94 -
working solutions was 5,230 parts per million, the largest eoneentration used in these experiments, the total concentration of hypoehlorous acid was decreased hy only 0.22$. It is assumed, therefore, that the effects noted in this study are not due to any shift in the chlorine hypoehlorous acid equilibrium caused hy the addition of chloride ion.
Effect of Chloride Concentration on the Reaction
Potassium chloride was added to the buffer solutions employed in the preceding experiments as a means for regulating the ionic strength values of the working solutions. This salt was selected for that purpose because the ionic mobilities of the ions resulting from solution of the salt in water are very nearly the same. Furthermore, the ions are affected neither by the oxidizing action of such strong oxidizing agents as chlorine nor the reducing action of such reducing agents as sulfides. However, certain of the results from the ionic strength experiments, which have been pointed out, suggest the strong possibility that chlorides affect the oxidation reaction in some manner, and it has already been mentioned that Higgins (42) noticed a stimulating effect produced by chlorides on the bleaching action of chlorine solutions. Consequently, experiments were planned with which to study this possible effect on the reaction, The procedure for these experiments has been described in the section concerning the experimental methods, and again at-


tention is called to the similarity between these experiments and those used to determine the effect of time* The pH and ionic strength values were held constant throughout the experiments at values of 7.0 and 0.1093, respectively, in order that the results of the experiments would he a comparable basis with the results of preceding ones.
The observed data from the experiments concerning the effect of chloride concentration on the reaction are presented in Table 20, and the calculated data are shown in Table 21* The calculated data relative to constants a and b of the empirical equation, R-s'at,' are briefly summarized for convenience as follows:
Chloride Concentration fnnm) Constant a Constant & fo Deviation of b from mean
0 2.01 0.0477 -5.2
385 2.07 0.0523 4*0
769 2.07 0.0464 -7.2
1150 2.07 0.0533 6,0
1540 2.16 0.0502 -0.2
1920 2.17 Q*sm. 3.6
Average 0.0503
Inspection of this summary shows that the full effect of the chloride ion on the reaction is apparently felt during the first minute of the reaction. This is indicated by the fact that there was a definite increase in the value of constant a with the addition of chlorides, whereas the


TABLE 20
EFFECT OF CHLORIDE CONCENTRATION ON THE RATIO OF CHLORINE TO SULFIDE REACTED* OBSERVED DATA*
Volume of buffer solution! 525 ml*
Sulfide added: 2,00 mg.
Concentration of sulfide solution: 2 nig/529 ml or 3.78 ppm. Chlorine added: 17.68 rag.
Temperature: 25 C..
Ionic strength: 0,1093
Chloride Reaction pH Chlorine Chlorine to Chlorine to
Concentration Time Reacted Sulfide Sulfide
(pom) (min) Ratio (nur) Ratio (mole)
0 1 6.98 9.17 4,59 2,08
t* 5 tt 9.58 4*79 2,17
.ti- 10 9.58 4.79 2.17:
ff 20 10.22 5.11 2.32
it 40 n 10,45 5.23 2.37
tt 80 n 11.07 5.54 2.51
385 1 6.97 9*27 2.10
5 a 9.98 4.99 2.26
ti- 10 tt 10.18 5.08 2.30
ff 20 w 10.64 5.32 2.41
ft 40 tt 10.89 5.45 2.47
tt 80 tt 11,80 5.90 2.67
Continued


TABLE 20 Continued
Chloride Reaction Chlorine Chlorine to Chlorine to
Concentration Time Reacted Sulfide Sulfide
.....,,-.....(#,}.............. (mp) Ratio (mg) Ratio (mole)
760 1 6.97 9.17 4,59 2,08
ft 5 0 05 4.98 2,26
ft 10 0> 02* 4.96 & # 20
ft 20 ft 10.63 5.32 2,41
w 40 n 10.81 5,41 2,45
80 11,03 5,52 2,50
1150 1 6.95 9.23 4,62 2.09
* 5 10,00 5,00 2,26
ft 10 it 10.15 5,08 2,30
ft 20 b 10.36 5.18 2.34
II 40 tt 11,20 5.60 2,54
It 80 11.84 5.92 2,68
1540 it tt 1 8,05 9.52 4,76 2,16
5 10,64 5.32 2.41
10 n 10,39 5.20 2.36
tt 20 it 10,98 5,40 2,49
it 40 ft 11.60 0.80 2.63
ft 80 11,88 5.94 2.69
1020 1 6.05 9.62 4.81 2,18
tt 5 tt 10,30 5,15 2.33
?t 10 10,84 5,42 2,46
ft 20 ft 11.16 5,58 13+03
ii 40 11.66 5 83 2.64
so '; tt 12.03 2.72


TABU 21
EFFECT OF CHLORIDE CONCENTRATION ON THE RATIO OF CHLORINE TO SULFIDE REACTED, CALCULATED DATA,
Source of data for calculations: Table 20, Basis for calculations: R at0, where
R Chlorine to sulfide ratio (mole), Reaction time (min),
a. and b are characteristic constants.
Chloride
Reaction
Constant
Constant
Calculated
$> Deviation
antration (ppra) Time (min) \ Chlorine to Sulfide Ratio (raole^ from Observed Ratio
0 1 2.01 0.0477 2.01 3.4
tr 5 tt tt 2.17 0.0
tt 10 it 2.24 3.2
* 20 tt ' 2. 32 0.0
40 tt : 2.40 1,3
tt SO tt ft 2.48 1*2
385 1 2.07 0.0523 2.07 -1,4
tt 5 tt- 2.25 0.4
tt 10 <- it : 2,34 1.7
ti 20 tt' 2.42 0.4
tt 40 it ". 2.51 i.e
ft 80 tt *.: 2.61 2.3
Continued


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