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The oxidation of sulfides by chlorine in dilute solutions

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Title:
The oxidation of sulfides by chlorine in dilute solutions
Series Title:
The oxidation of sulfides by chlorine in dilute solutions.
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Goodson, James Brown,
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Gainesville FL
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University of Florida
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English

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Chlorides ( jstor )
Chlorine ( jstor )
Mathematical constants ( jstor )
Nitrogen ( jstor )
Oxidation ( jstor )
pH ( jstor )
Potassium ( jstor )
Sulfates ( jstor )
Sulfides ( jstor )
Sulfur ( jstor )

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University of Florida
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University of Florida
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Copyright James Brown Goodson. Permission granted to the University of Florida to digitize, archive and distribute this item for non-profit research and educational purposes. Any reuse of this item in excess of fair use or other copyright exemptions requires permission of the copyright holder.
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021711075 ( alephbibnum )
13231918 ( oclc )

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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


THE OXIDATION OF SULFIDES BY CHLORINE

IN DILUTE AQUEOUS SOLUTIONS












By
JAMES BROWN GOODSON, JR.










TABLE OF CONTENTS

Page
INTRODUCTION AND REVIEW OF THE LITERATURE .�.�. 1
STATEMENT OF PROBLEM o. oooooo.ooooo 13

EXPERIMENTAL NETHODS .o. 14
Preparation and Standardization of Solutions . . 14
Description of Apparatus 23

Experimental Procedures sooe*~*o. o .o. *. o 0
EXPERIMETAL DATA . 42
Precision of Determinations .�. 44
Effect of Concentration on the Reaction ,.o 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
$U4AY0004# 060g*0 b0 4. 4 44O49400 0**@O 0 .aJl4, tQ00 08,,*
Mh y 1 08

BIBLIOGRAPHY U#*000*404**00000409011 ACKNOWLEDGEMMNTS -0. . 115
BIOGRAPHICAL ITEMS *o.o0.0 .o.o �.�. 116
COMMITTEE REPORT � , 117










INTRODUCTION AND REVIEW OF 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 are not uncommon. The origin lof 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 hasanught 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, sulfiden-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 sulf ides 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


-wI -







as2 40


method part of the sulfide, the amount being largely depend-. ent upon the hydrogen-ion concentration of the water, is expalled to the atmosphere as hydrogen svulfide, and the remainder is oxidized to some degree by the dissolved oxygen in the aerated water. The method is attended by many difficulties or 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.
Chlorination as a means for sulfide removal from natural waters is of rather recent origin, It appears to date from 1928, (1) in vhich year it was resorted to at Beverly Hills.,

California, after aeration alone had been proven to be in-. effective 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 statement: *Mention should be made, however,_to the practice at Holland, New 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. Rideal, (3) in reporting the results of a long








-3-


series of experiments dealing with hypochlorite treatment of septic sewage liquors in the 1908 report of the Royal Commission 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 bleach 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: H28 S Cl2: 2HC1 S 8. He concludes that for the removal of each pound of hydrogen sulfide 2.1 pounds of chlorine are required. Pomeroy and Bovine (5) have the following to sayt .When 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: Cl2 / H2S S ,t 2HC1.o. 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 the chlorine dosages Just sufficient







.4 g


to eliminate the sulfides from various samples of septic sewage found that the ratio of chlorine to sulf ides 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 or the chlorine. Experiments by Powell and von Lossberg (6) Indicate that the chlorine dosage for removal or hydrogen sulfide from some natural waters in which they were Interested approximates the etoichiometric value dictated by the following reaction: 4C12 41120 1128 : 112804 / 81C1 The 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. How-. ever, 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 oxidizing potentials of the oxidizing










couples used. In the following discussion couples with oxidation potentials in the vicinity of or greater than one volt are considered to be strong oxidants, end those with potentials less than one volt are considered to be weak oxidants. Among the weakest oxidants studied were neutral to alkaline solutions or suspensions of potassium chromate (9,10), potassium dichromate (11), ammonium chromate (12), lead chromate (13), silver chromate (13), mercurous chromate
(14), barium chromate (15) and solutions of various water soluble aromatic nitro compounds (16)o The principal products of the oxidation of alkali sulfide and hydrogen sulfide by these oxidants were found to be polysulfides, free sulfur, thiosulfates, 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, thiosulfates and polysulfides, to a secondary reaction between

*active* sulfur and hydroxyl-ion, which they write as follows: 60H / 12S u %03sI./ W5- / 3H 0. 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-


-05-4








-6.


gen sulfide by iodine, which is considered to be a weak oxidant. That oxidation proceeds quantitatively to free sulfur in an acid medium (17) 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 trioxide (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 case it was found that under certain conditions relative to concentrations a dithionate

occurs as a final product.
The preceding review of the work of some of the investi-







-7 -


gators who have 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 or sulfides and may be the only end-product In acid solutions when weak oxidants are employed (18, 27). In basic. solutions there are changes noted In the final products of the oxidation (17, 18, 20). These changes are evidently ex-. plained 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 effect of the hydrogen-ion concentration on certain oxidants Is illustrated in the specific ease where nitrous aid was employed as the orwF4vting agent.
Mention was made by Bagoter (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







go8a-


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 oxidizing characteristics is noted in the case of iodine (17). Although the oxidation potential of the Iodideiodine couple Is not dependent upon the pH or the solution at pH values below 8, at higher pH values the iodine will react with the hydroxyl.-Ion to yield hypoiodite (30) thereby alterIns the or141vins 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 oxidized 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: 6011 / 88 28~ S 203" / 3110 An excess of sulfur was found to

yield pentasulfide by a secondary reaction. As pointed out by Bullock and Forbes (16) the reactions do not proceed rapidly at 250 C. with ordinary rhomb ic 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 thiosulfates in alkaline solutions (33, 34).

Now let us consider some of the properties of chlorine as an oxidizing 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 chloride-hypochlorite 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. Higgins (35, 36), Rideal and Evans (37) and Remington and Trimble (38) have studied the effect of acids on the oxidizing properties of hypochlorite solutions, It is apparent from the works of these investigators that the oxidizing power-of hypochlorite solutions can be markedly increased by the addition of weak acids. Rlideal (39) suggests that free hypochlorous acid is the effective oxidizer 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 hypochlorous and carbonic acids to show that free carbonic acid will liberate free hypochlorous acid from hypochlorite solutions. Higgins (36) 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. He also attributes the oxidizing properties of such solutions to the active mass of free hypochlorous acid and explains that the boric acid liberates free hypochlorous acid while the hydrochloric acid liberates free chlorine. Veiss (40) is discussing the kinetics of chlorine bleaching claims that the active agent is chlorine monoxide or undissociated hypochlorous acid. He mentions another interesting point in connection with the oxidizing 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 Higgins (42) that chlorides have an accelerating effect on 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-







-- 11 -


rides were not present, Chlorides produced by the reduction or the hypoehiorites during the bleaching reaction appear to have a negligible accelerating effect. It is mentioned by Higgins (43) that there is a secondary reaction between hypochiorous acid and neutral chloride whereby nascent chlorine or 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. Perellman and Lselyakina (45) found that hydrogen sulfide in acetylene gas is quantitatively oxidized to sulfate when passed through a solution of sodium hypochlorite and propose a quantitative method for the determination of hydrogen sulfide in acetylene based on this reaction. Some comprehensive work has been done by Choppin and Paulkenberry

(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 faot that the end-products of the oxidation are sulfur and sulfate, the ratio depending upon such factors as relative concentrations of the original reactants,, hydrogen-ion concentration of the reaction medium,







- 12 -


temperature, standing time and rate of addition or reactants. 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 or these factors on the secondary reactions,'

For example, higher temperatures were found to increase the proportion of sulfate, and p1* 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, 6WU / as 28
3
S 203- / 31*20, where Increased temperatures and high alkalinities favor the solution or sulfur (32, 18). The products are subsequently oxidized 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 pH 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.










STAThONgT OF PROBISEH


*The object of this investigation was to make a study

of the oxidation of' sulfides in very dilute aqueous solutions, such 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 reactants, 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.


- 13 -









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 iodometric or todimetric 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.

Prparatgn and Standard@ gion of Solutips

Standard gg I ppt~sLum dichrompp aqlutto~. Some reagent grade potassium dichromate was pulverized in an agate mortar and dried in an oven at 1500 - 2000 C 9.808 grams of the material were dissolved in distilled water and diluted exactly to 2 liters,
Standard g g potasspm dchrompte solut rn. 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 -


Suff'i acid. Concentrated reagent grade sulfuric acid was used where this acid was called for in the procedures.

Jotaqxsum iodide. Reagent grade potassium iodide
crystals that have been tested for the absence of Iodate were employed in the investigation.
Starch ind~~9,r soution, 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.
Z acetae ,lu~ qn. 240 grams of reagent grade zinc

acetate were dissolved in one liter of distilled water.

0&01 AIidine qqjkption� 2,54 grams of reagent grade iodine were dissolved in several ml. of water containing 8 grams of iodate-free potassium iodide. This solution was diluted to approximately 2 liters. No standardization was necessary.
,ydroch orianaoi. Concentrated reagent grade hydrochloric acid was used where this acid was called for in the procedure.
Acetic agid. 500 ml. of glacial acetic acid was diluted to one liter with distilled water.







- 16 -


$1 sod$M thosWlfpty fqutpn. - Approximately 200
grams of C. P. sodium thiosulfate 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 thiosulfate solution, starch being used as the indicator.

PAF. sodium this atq gqutgon. This reagent was prepared by diluting a measured amount of the aged and standardized 0.1 N sodium thiosulfate 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 thiosulfate, starch being used as the indicator.







- 17 -


Chggrine agg. 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 mlliliter. When the atmosphere was excluded from the solution by a simple cheek 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 daill 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 thiosulfate with starch as the indicator.
gStock ,uftde 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 Na2S9HNO 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 standardized daill 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 zinc acetate solution, and the resulting mixture was diluted with 110 ml. of distilled water. 50.00 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 standardized 0.01 N sodium thiosulfate, 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 45 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 fg pqutions. Clark and Lubs buffer
solutions having pH values of 5.00, 7,00 and 9,60 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 buttffer 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 cabe 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 ease of the pH 5 buffer solution the limited solubility of the buffer materials interfered with this plan, and a con-







- 20 -


centrated 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 Lubs series (49), with potassium chloride being added in addition to equalize the tonic 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 standardized, 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. Hoever,







TABLE 1
COMPOSITION OF CONCENTRATED BUPTR SOLUTIONS AND CHARACTERISTICS OF WORKING BUPPER SOLUTIONS. pH of Components of concentrated buffer %1/525 ml Characteristics of
solutions
working (gag/tq) working vo'kpg buffer soiytipnq
buffer KHC84H 4 KHPOP H 3BO NaOH KC1 buffer *I without Added C01 I witk solution. add9d KC (pvm) added KC

5.0 107.2 . *---- 10,02. 94.59 50 0.04885 2140 0.1093
6,0 214.4 **--- 7.184 26,87 25 009210 808 0.1093
8.4 * 142.9 ----- 10.58 53.33 25 0,07520 1208 041093
6., -- 142.9 ** 14.95 37.05 25 0.08560 840 0.1093
6.8 --- 142.9 - 19.70 19.36 25 0.09690 438 0.1093
7.0 71.45 ---* 12.44 -- . 25 0.05463 0.0546Z
7.2 --- 114.4 ---. 23.52 20.77 25 0.09600 470 0.1093
7.4 114.4 --- 26.54 9.492 25 0.1032 215 0.1093
8.0 214.4 . 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 onlorine losses In the blank determinations with this buffer solution were of' the same order of' magnitude as those with the phosphate and boric acid solutions. Therefore, It Is concluded that no error due to such a chlorination has been introduced'.

Pqtassiwn chloride solpztignso 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 97.88 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 7.6) in a series of' working buffer solutions originally containing 50 ml, of concentrated buffer per 525 MI. of solution would yield a series of working butter solutions having a constant ionic strength value of' 0,1093 and chloride concentrations varying by equal incremnts from zero to approximately 2,000 parts per million. Such a solution was calculated to contain,88.52 grams of potassium chloride per liter.







- 23 -


These potassium chloride solutions were prepared In nitrogen.-filled volumetric 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,

DesqOptitir of Apra"tug

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 and-on a line passing through the center, They were fitted with rubber stoppers that had been carefully ground to the proper size to insure a positive closure, The gaskets were out 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, Vhen 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 2-1iter, 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













F/ G URE
SCHEMATIC DIAGRAM OF NITROGEN ASSEMBLY


Legend
A - S3eel nilroyen cylinder B- Need/e a/vale w1/h pressure gauge C - Pressure reyu/la/or D- ol/ufion storage bo/l/e E - Nifrogen 7 an//iold F- Reaction vessel
G- Rack for supporting reaction vesse/s
H- / Lifer beaker


to s a







- 26 -


sleeves protruded. These sleeves were for the admittance of pipets so that .the solutions could be withdrawn while a stream of nitrogen kept back the air, The manifold (E) was constructed from 2.5 em. tubing about five feet in length, and it was fitted with nine, equally spaced, 0.5 oa, 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. vessel by the .nitrogen during the filling process.

Shrown in Pigure 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 pipet (B) and a return tube from the top of the pipet, This return tube closed the system and prevented exposure of the diluting water to air. The pipet (B) was a 500 ml., semi-automatic, volumetri pipet especially designed







o 27 t


FIGURE 2
SCHEMATIC D/A GRAM OF APPARATU61S 1/SED //V
PREPARE AT7/0N OF WORKING SOLLITiONs


Legend
A - Di/u/inyg wyoer s/oroae tfnk
2 - 500 m/, semi- ou/ona//c, ro/amnetrl pipe C - Stock s/fid'e ao/u// storope f/sk D - / m/ mi/cro bure/ f -- '//rojen man/'fo/d
- - Reac/bn vesse/ S-- Ch/ornC waler slorae bo///e / -- 50 m/ - bure? _I - /0 m/ /-/'cro buref







- 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 (E) to deliver 'a stream of nitrogen into the vessel and thus keep back the air. The stock sulfide solution was kept in a 2-liter, erlenmeyer flask-'(C) that was fitted with a 3-hole rubber stopper. A nitrogen tube from the manifold (E), a stoppered vent tube and a siphon to buret (D) passed through the. three holes. The buret (D) was a 10-mi,. micro buret to which a side arm had been added just above the stopcock, A nitrogen filled, rubber bladder of the type used in an Orsat gas analyser closed the top of the buret and prevented the exposure of the sulfide solution to air. A 10 liter, black, enameled bottle (G) with an outlet near its bottom served as storage for the chlorine water. The bottle was stoppered with a 1 hole rubber stopper through which protruded a simple check valve that acted as a barrier to the diffusion of gases to and from the atmosphere. The check valve consisted of a short length of glass tubing stoppered at one end with a 1-hole rubber stopper through which was inserted a short section of smaller diameter glass tubing that had a rubber policeman stretched over its lower end. A slit in the policeman-allowed air to enter the bottle only when solution was being removed, The chlorine water storage










bottle van connected to a 50 ml. buret (H) and a 10 ml. buret (1) 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 pinchclamp,
The thermostat used In this Investigation was of the conventional, manually operated type, The water bath conslated 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 thedimensions of the box allowed for a volume of waterthat 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* Tempera. ture 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 lea bath. The proportion of water passing through the cooling coil 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. le C.
The Beckman pH meter., model G,, van used for all pff

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.


gyper aent Prpcadurye

procedure for preparation o wqr4ng p9qtqgn. This
procedure was complicated by the necessity for keeping oxygen out of contact with the solutions and the various reagents as the solutions'vwere 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 stoppereds 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 vesselvessels 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, Fiigure 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 pipet 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 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 burst (see D, Pigure 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 coleuc lated from the daily standardization 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 size of the buret de. pending 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 standardization 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,

Snagt~pcaj 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 40


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 or 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 iodometric method for the determination of chlorine in-water (50). The reaction in the vessel was stopped at the end or the reaction time by the addition of 0,75 grams of potassium iodide dissolved in enough acetic scid solution to lower the p11 or the working solution to a value between 3 and 4, Usually 10 ml. of the acid was suf. ficient. -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 or distilled water from a wash bottle. The liberated Iodine was titrated with the standardized 0.01 N sodium thiosulfate solution. using 5 ml, of starch solution as the indicator# 'The quantity or residual chlorine was calculated from the titration.
The method used for the quantitative determination of

residual sulfide was based on the standard iodimetric 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 N iodine was measured Into the reaction vessel by 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 standardized 0,01 N sodium thiosulfate 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 thiosulfate solution. The amount of residual sulfide was determined from the amount of iodine that reacts.R x - mtjF j detpri19 the e feet o_ cqnceqtrtns. 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 19, 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 a 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 tht 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 con-. taining the sulfide solutions in each experiment were 2, 4, 6, 8 and 10 times the sulfide dosage for he 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 9.0. The reaction time and the temperature were held, constant throughout the experiments at 20 minutes and 250 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 pipet 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.

In phes exerimep dnten rece ign g a sedts o n 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 25o 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.

dperima3tgnI&eqteru4np the extent athir4 9 IA md*t
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-mle 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 persist. ed 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 250 C. for all runs's The deviation that has been mentioned concerning the prepa-. ration of the p11 5 working buffer solution occurred again during the se experiments.
)Sxperime1t ~ eermine tha effect or tjg The concen.tration 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 milligram 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 chlorine 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 I,- 5,-10, 20, 40 and 80







- 39 -


minutes, respectively. The sme 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 experiments, 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 experients was held constant at 250 C. The deviation occurring in the preparation of. the pfH 5 working buffer solution is again pointed out in connection with these experiments,

J.periments &t detrmin9 A Affg. tem " prtpturq. Experiments identical to those described in the above paragraph were. run at pH values 5.0 6, .0, 7.0, 80 and 9.0, but temperatures of 150 and 200 C, were substituted at each pH value for the 250 C. employed in the previous runs.

aer",Rtp 11 detraniVi9 ID &eg.c1 Ig ionicg stpth. These experiments wcre 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 perm t. 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.
!yperpints j detprmau th e g chl~rAdps. The chloride experiments were run in the 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 250 0C. 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 addeds 50 ml. portions of concentrated buffer solution (pH 7.0) were added









to 475 ml. portions of diluting water obtained by withdrawng 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.


congentratqn. 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.










EXPERDWRTAL DATA


The results of the various experiments conducted during this investigation are expressed as ratios that show the number of units of chlorine 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 realize 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 endproducts 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.:
C12 / Sa2Cl- / S

4C12 / S__ t 4120: 8=0 St S04"',
In accordance with these relationships the respective ratios for the formation of free sulfur anid sulfate are the following:

Mo,]ee, of c9)prine :1
Mole of sulfide

MOORe Or c ,r1:WOe
Mole of sulfide
-42-







- 43-4


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 determin

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 oxidized. 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 res pect to sulfide) and

X 4Y Z 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 oxidized expressed In the same units.










Similarly, the amount of sulfate can be calculated by multiplying 2.99Y by the quantity of sulfide oxidized 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*

~rec*Vion pf Dqterj~n tAons

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 which 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 :


1.00 1.10
1.20 1.30
1.40 1.50 1.60 1 70 1.80 1.90
2.00 2.10
2.20 2.30
2.40 2.50 2.60 2.70 2.80 2.90 3.00 3.10
3,20 3.30
3.40 3.50 3.60 3.70
3.80 3.90 4.00


Moles Sulfur
per Mole of anrlYwa t1


1.00 0.967
0.934 0.900
0.867 0.833 0.800
0.767 0.733 0.700 0.667 0.633
0.600 0.567
0.533 0.500
0.467 0,433
0.400 0.367 0.333 0.300
0.267 0.233 0.200
0.167 0.133
0.100 0.067
0.033 0.000


Moles Sulfate per Mole of 1 a1sr - Irl


0.000
0.033 0.066 0.100 0.133 0.167
0.200 0.233
0.267 0.300 0.333
0,367
0.400 0.433
0.467 0.500
0.533 0.567 0.600 0.633
0.667
0.700 0.733 0.767
0.800 0.833 0.867
0.900 0.933
0.967
1400


MgIep Sulfur Mole Sulfate
[vi','


29.3 14.2 9.00 8.52 5.00 4;00
3.30 2.75 2.33
2.00 1.73 1.50 1.31
1.14
1.00
0.876 0.763
0.667 0.579 0.500 0.428 0.364
0.304
0.250 0.200 0.153 0.111 0.071
0.034
0.000


- 45 -


UUtb~nLU~7 ~~~ U~CL~iU~J X1I )4(*11










TABLE 3

P#ISCIS:ION OF )DETEMINATIONS AT VARIOUS pH VALUESVolume of buffer solution 525 .ml.
Sulfide added: 2.00 mg
Concentration of sulfide solution: 2 mg/529 ml or 3.78 ppm. Chloriie added 17.68 mg.
Reaction time: 20 min,.
Temperature: 250 C,


7 *
7.


- .
7.


Il Chlorine
Reacted
(mg)
93 15,60

" 15.55

15.58


89 12.20

12.09 12,17


05 10.44

10.10

10.22


92 8.35

* 8.51

8.35

" 8.24


Chlorine to Ch Sulfide


7 *80

7.78 7.79


6*10

605 6.09


5.22 5,05 5.11

_ _ .3 . .
4.18 4.26

4.18 4.12


lorine to % Deviation Sulfide from
to (mole) *vrage

3.53 013

3.52 0.0

3.52 0.0


2.76 0.4

2.74 -0.4

2.75 0.0


2.36 1.7

2.28 -1.7

2.31 -0,4

_. -. . . . ~.
1.89 0.0

1.93 2.1

1.89 0.0

1.86 -1.6


- 46 -


4.


. I
5


.







- 47-


The experiments concerning the effect of the concentrations of the reactants on the oxidation were designed to in. dicate the extent or 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 9.0 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 SILFurIJ REACTED.
SULFIDE CONCENTRATION IS 1.90 PPM.


Volume of buffer solution Sulfide added:2 Concentration of sulfide s Reaction time: Temperature


525 ml.
1.00 g, solution 1 g/527 ml or 1.90 ppm.
20 min.
250 C.


pH Chlorine
Added
(mi)


2.00 4.00 8.00 8.00
10.00

2.00 4.00 6.00 8.00 10.00

2.00 4.00
6.00 8.00 10.00

2,00 4.00 8.00
8,00 10.00

2.00 4.00 6.00 8.00 10,00


Chlorine Reacted (Ia


2,00 4.00 6.00
7.42 7.91

2.00 4.00 5.06 5.56 6.00

2.00 4.00 4.82 4.76
4.93

2.00

4.15 4.12
4.21

2.00


4.23 4.31


Sulfide Reacted ( mr


0.318
0,574 0.826
1,00 1.00

0.470 0.848 1.00
1.00 1.00

0.510 0.952 1.00 1.00

0.592
1.00 1.00 1.00 1.00

0.562
1*00 1.00
1,00
1.00


Chlorine to Sulfide


8628 6.97 7,26
7.42 7.91

4.25
4.72 5 06 5.56 6.00

3,92 4.20
4482 4.786
4.93

3.38

4.15
4.12 4.21

3.56
---

---






4.23 4.31


Chlorine to Sulfide
R (ii~


2,84 3.16 3,29
3.36 3.58

1.93
2.14 2.30
2.52 2.72

1.78 1.90
2.18
2.16 2.23

1453

1.88 1,87 1.91

1.61

1,92 1.95


- 48 -


4,88




5.89

ft


7.05

ft


7.92
a




8.95
a
a
# ft


IgI-* 7g4r, pg votV %g tgo










TABLE 5


EFFECT OF CONCENTRATIONS ON THE RATIO OF CHLORINE
TO SULFIDE REACTED.
SULFIDE CONCENTRATION IS 3.78 PPM.


Volume ofrbuffer solution:


Sulfide added: Concentration of sulfide
Reaction time: Temperature:


pH Chlorine
Added -n (M)I


-Chlorine Reacted
1 *y


4.00 8.00
12.00 14,98 168041

4.00 8.00
11.04
12.10 12.80

4.00 8. 00 9,59

10.34

4.00 8.14,
8,1 8.15


4.00

7.50 8.23 8.27


solution 2 mg/529 ml


Sulfide Reacted
Im %


0.694 1.18 1.70 2.00
2,00
0.901 1.61
2.00 2.00
2.00

1 *89 2,00 2.00
2.00

1.18
2.00
2.00 2.00 2.00


4.90
N
I
I
I

5.89
I
I
I
*
a
a
7.05
M
I
V
I

7.92

.w
V
N
N
S
8.*95
N i,
I
U


Chlorine to
Sulfide Ratio (mr)


5.77
6.78 7.06 7.49 8.02

4,44 4.97
5.52
805 6.40

3,92 4.24
4.80 4.98 5.17

3.39

4.07 4.08 4.14

3.54

3.75
4.12 4.14


525 ml.
2.00 mg., or 3.78 ppm.
20 min, 250 C.

Chlorine to
Sulfide
Ratio (mole


2.62
3.07 3.20 3.39 3.63
2.01 2.25
2.50 2.474
2,90

1.78
1,92 2.18 2.26
2.34

1.54

1,84
1.85 1,87
1.60

1.70
1.87 1.87


- 49 -


F 91


4.00 8.00
12.00 16,00
20.00

4.00
8,00
12.00 16.00 20.00

4,00
8.00
12.00 186.00
20.,00

4.00 8.00 12.00 16.00
20.00

4.00 8,00
12,00 18.00
20.00


1.13
2.00 2.00 2.00 2.00


Z~C -~r-r hP~~-?C


"""" """ "










TABLE 6


EFFECT OF CONCENTRATIONS ON THE RATIO OF CHLORINE
TO SULFIDE REACTED.
SULFIDE CONCENTRATION IS 5.65 PPM,


Volume of buffer solution: Sulfide added: Concentration of sulfide solution: 3 mg/531 ml Reaction time: Temperature:


pH Chlorine
Added
(mr)r


Chlorine Renactid
fme%


6,00
12.0) 18 00 22, 82 24.12

6.00
12.00 16.91 17.72
18 30

6.00
12.00
14.15 14.85 15,61.

6.00

12.06 12.26 12.65

6.00

11.94 12.35 12,66


Sulfide
Reacted
mr)


1.01
1,74 2.51
3.00 3.00

1.27
2.29 3.00
3.00 3.,00

1.50
2.76 3.00 3.00 3.00

1,74
3.00 3.00 3.00 3.00

1.66
3.00 3.00 3.00 3.00


Chlorine to
Sulfide nftfin (mv/)


5494 6.90 7.17 7.61
8,04

4.72
5.24 5.64 5.91 6.10

4.00
4.34 4.72
4.95 5.20

3.45

4.02
4.09 4.22

3.62

3.98
4.12 4.22


525 nml.
3;00 mg, or 5.65 ppm.
20 min.
250 C.

Chlorine to
SSulfide Ratio (mole)


2.69
3.12 3,25 3.45
3.64

2.14 2.37 2,55 2.68 2.76

1.81 1.97 2.14
2,24 2.36

1.56

1.82 1*85 1.91

1.64

14,80
1.87 1.91


- 50 .


4.90 it
a
a

5.89
p
a
a

7.08
U
U


7,92
a pr


It
8,95

a


6 00 12,00 18,00
24.00 30,00

6,00
1200 18.00
24,00 30,00

6.00 12400 18.00
24.00 30,00

6.00 12.00 18.00
24.00 30.00

6.00
12.00 18.00 24.00 30.00


Ratio (molel


1, F'v T- . F -N"7 . -T- -7 - N117 N


.







- 51 .


quired for the more concentrated solutions. At any rate the results for the 3.78 and 5065 part per million solutions are summarized 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 clearcut dependence of the experimental ratio upon the hydrogenIon concentration of the reaction medium, The ratios were quite large In a reaction medium having a piT value of 4.90., Indicating that the larger portion of the oxidized sulfide was oxidized all the way to sulfate., However, the ratios docreased as the pHT value of the reaction medium was raised, and it is apparent that limiting values were approached in alkaline media at a pff 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 oases with an increase in the ratio of chlorine added, This increase was quite substantial in acid solutions, but it became less marked as the pHT 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


EECT OF CONCENTRATIONS ON THE RATIO OF
TO SULFIDE REACTED.
SUMMARY OF EXPERIMENTS


pH Mole Ratio of Chlorine
to Sulfide Reacted
Sulfide Cone. Sulfide Cone.
3.78 ppm 5.65 ppm
Table 5) (Table 0)


2.62 3.07 3.20 3.39 3.03

2.01
2.25 2.50 2.74 2.90

1.78 1.92 2.18
2.26 2.34

1.54

1.84 1,85 1.87

1, 60

1.70
187 1.87


2.69: 3.12
3.25, 3.45 3.64

2,14 2.37
2.55 2.68 2.76

1.81 1.97
2.14
2.24 2.36

1.56

1.82 1.85 1.901

1.64

1.80 1.87 1.91


Average
Ratio


CHLORINE


% Deviation
from
Average


2 66 3.10'
3.23 3.42 3i64

2.08 2.31
2.53 2.71
2.83

1,80
1.95 2,16 2.25
2.35

1.55

1.83 1.85 1.89

1.62

1.75
1.87 1.89


1.3
0.8 0.8 0.9
0.2

3.1
2.6 1.0 1.1 2.5

0.9
1.3 1.0 0.5 0.5

0.7

0.6 0.0
1.1

1.3

2.8 0.0
1.1


- 52 -


4.90


I

5 *89

I


I
U
U ,8
a





7.92
*
p
*



U

8.95
U
U






I


J-~ -r 1' -ra cir








- 53 -


FI G LRE 3 EFFECT OF CON CENTRATIONS ON THE RATIO OF CHLOR/NE
TO SULFIDE REACTED.
A AVERAGE RATI/0 FOR 3.78 v, 5.65 PFM SULFIDE S30LTIONS PLOTTED.

(See Table 7)



3.70


3.50-H


" 3.303.10- O




2.70oN.0

P 2.30o
2.10
SO "p/1= 7.92(0)
p/9- / 8.95 (A)




l5D I
0 2 4 6 8 10
Ro//o of Ch/or/ne Added to Su//ide - My.







- 54 *a


Figure 3 capably illustrates another important consideration, It is noted from Tables 4 8 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 unoxidised 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 chlorine 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 unoxidixed sulfide in reaction media of various pH values lie between the following values

PH Chlorine Dosage


4,90 688

5.89 4 6
7.05 4 -6
7,92 2 -4
8.95 2 -4







i-55-


EffROt Or' Tioe thjeato

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. 7his Initial portion of the oxidation was found to give way to a much slower process. The experiments shoved 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
(48), it was decided that the study of the oxidation rate could best be made In two parts,, Vie 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 standardized chlorine 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 nre 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


EXTINT OF THE TM~MITATE REACTION BETWEEN CHLORINE AND SULFIDE.

Volume of buffer solution 525 ml.
Sulfide added: 4.00 mg.
Concentration of sulfide solution: 4 mg/533 ml or 7.51 ppm. Temperature: 250 C.
Reaction time Sulfide solution is titrated with: hlorine
water, using starch and KI as the indicator.


pH- Chlorine Reacted

4.93 9.34
" 9.31
* 9.48
* 9.47


5489




7.05
a










8.95
*
S



792
p
*1

S
8.95
p
'I


8,86
8.87 8,89 8.89

9.62 9.77
9.83 9.84

11.17 11.07 11.13
11.23

12443 12.24
12.10


Chlorine to
Sulfide
Ratio (mrl


2.34 2.33 2.37 2.37


2.22 2,22 2.22 2.22

2.41 2.44 2.46 2.46

2.79 2.77 2.78 2.81


3.11 3.06 3.03


Chlorine to
Sulfide
Ratio (manlfa


Average


Average


1.06 1.06 1.07

1.065

1.01 1.01 1.01


1.09 1.10 1.11

Average 1.10

1.26 1.25 1.26
128
Average 1.26

1.41 1039

Average 1,39


%Deviation
from


-0,5 0.5
0.5

0.0 0.0
0.0 0.0

0.9 0.0
0.9 0.9


0.0
-0,8 0.0 0.8

1.4
0.0
-1.4


- 57 -


-ri---7 ---'-r~ ~7TI ~-"~







- 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 mwaters, 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 oxidize 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 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 & and . 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 ,. 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 k 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


Volume of buffer solution: Sulfide added: Concentration of sulfide s Chlorine added: Temperature:


525 rml
2.00 nag. olutiont 2 mg/529 ml or 3.78 ppm.
17.68 mg. 250 C.


pH Reaction
th e ( b)}


4.90

it


5.89


I,$
a
a



i
7.05
w


it
w





It it


79
it
8 *95 it 'V it it it


Chlorine
Reacted % -t I


14.92 15,05
15.27 15.30
15.46 15,44

11.38 11.92 11.96 12.20
12.38 12.50

9.72 9.60
10.24 10.46
10.87 11.22

8.24 8.39
8.45 8.38 8.75 8.70

7,60
7.99 8.09 8.26 8,41
8.53


Chlorine to
Sulfide nern4 &% f1".r


7.46
7.53
7,64 7.65 7.73 7,72

5.69
5.96 5.98 6,10
6.19 6.25

4.86
4.80 5.12
5.23
5.44
5,61

4,12 4.20
4.23 4.19
4.38 4.35

3.80 4.00
4.05
4.13 4.21 4.27


Chlorine to
Sulfide Rat~m uib1


3.38
3.41
3.46 3.47
3.50 3.50

2.58 2.70

2,76 2.80 2,83

2,20
2.18 2.32 2.37 2.46
2.54

1.87 1,90 1.92
1.90 1.99
1,.97

1.72 1.81 1.83 1.87 1.91 1.93


Logarithm of Ratio IM


0.529 0.533 0.539
0.540
0.544 0.544

0.412
0.431 0.433 0.441 0.447
0.452

0.342 0,339 0.366
0.375 0.391
0,405

0.272 0.279 0.283 0.279 0.299 0.295

0.236 0.258 0.263 0.272 0.281 0.286


- 60 -


We Y' "W!g _7Y g 4gy








- 61 -


FIGURE 4
EFFECT OF TIME ON THE RATIO OF CHLORINE TO SULFIDE REACTED
(See Table 9 )


0,4 5 ---- (3



0.5e


N0




0.45
















0.30 - 0
04C







-N O
SpH: 7.9

0.30 -


O 20


.2 o. 0.6 0.8 /.9 1.2
Lo9ari/hm of Time -AMia,











TAIBE 10


EFFECT OF TDE ON THE RATIO OF CHLORINE TO SULMfDE REACTED.
CALCULATED DATA

Source of data for calculations: Table 9.
Basis for calculations R = atb , where.
R : Chlorine to sulfide ratio (mole),
: Reaction time (min),
a and b are characteristic constants.


pH Reaction
Time
(min)


4,90
a
n
a



5 .89
#
a
a
a

7.05 ft




7.92



ft


8 *95
#
ft
a ft

'I6) ft


Constant Constant
. 1 .1_ h


3.39
-t
U ft
a ,t ft

260
ft f"

frt

a
2.04
a
a



1*87

w



1.74
ft ft ft ft ft


0.00720
a a a a a

00193
ft






0.0497
S




ft


0.0124
0 ft







0.0497
ft ft ft
ft






ft
0.0124.
fit




0.02431
ft U ft p


Calculated Chlorine to
Sulfide
Ratio (molt


3.48

3.50


2.68 2,72
2.76 2.79 2.83


2.21 2.29 2.37
2.45 2.54


1.90 1,92

1.94 1*95 1,97

1.74 1.80 1 84

1.90 1 .9


% Deviation from Observed
Ratio


0.3 0.6

-0,30
-0., 6 0.0

0,8
-0,7 0.4 0.0
-0.4
0.0

-74;5 1.4
-1,3 00
-0.4
0.0

0,0
0,0

2.1
-2.0
0.0

1.2
-0,6 0,95
0,0
-0.5
0.0l


-62 -


7~









ments planned to illustrate the effect of temperature on the oxidation extends from 150 to 250 centigrade. This range was chosen with the thought that it includes the temperatures of the great majority of sulfide-bearing -aters. Purither more, these temperatures were found to be 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 150 or in excess of 25o were employed. Also to be considered was the probability_.tat temperatures much greater than 25 degreesV 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 150 and 200 C., and the results were cmpareo with the 25 degree results.
The data observed during the 15 and 20 degree experiments are presented in Tables 11 and 12, respectively, .and t3ey are summarized 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 th variations in temperature. Indeed, in the large majority of cases the maximum deviations between the ratios for the varl-










TABLE 11


EFPBCT OF TEMPERATURE ON THE RATIO OF CHLORINE TO
SULFIDE REACTED.
OBSERVED DATA


Volume of buffer solution Sulfide added: Concentration of sulfide solution: 2 mg/529 ml Chlorine added: Temperature:


525 ml.
2.00 mg. or 3.78 ppm;
17.68 mg.
150 C.,


Reaction
Time (Mii I


1
5 10
20 40 80

1
5 10
20 40 80

1
5
10
20 40 80

1
5 10
20 40 80

1
5
10
20 40 80


Chlorine Reacted
(m I


14.76 15.17
15.24
15.48 15.56 15.68

11.36 12.05
12.02 12.10 12.38
12.43

9.41 9.93 10.35 10.36 10.57
11.21

8.14 8.23 8.52
8.42
8.64 8.90

7.06 7.76
7.95 7.99 8.17 8 33


Chlorine to Sulfide
Rati (m I


ga f v


7.38
7.59 7.62 7.74
7.78 7.84

5.68 .6.03
6.01 6.05 6.19 6.21

4.71
4.97 5.18
5.18 5.29 5.61

4.07 4.12 4.26 4.21 4.32
4.45

3.53 3.88 3.98 4.00 4.09
4A .


Chlorine to Sulfide Ratiqu (mole)


3.34 3.44 3.45 3.50 3.52 3.55

2.57 2.73 2.72
2.74 2.80
2.82

2.14 2.20
2.34 2.34 2.40
2.54

1,84 1.87 1.93 1.91
1.96
2.02

1.60 1.76 1.80 1.81 1.85
1.89


-64.-


4.93
N
N
N


5.97
I,
N

5* 9
N
N



7.14
U

*



8.02
N
U



9* 1 9*15

N
N
N


1,89










TABLE 12


EFFECT OF MOSPERATURE ON THE RATIO OF CHLORINE TO
SULFIDE REACTED. OBSERVED DATA


Volume of buffer solution: Sulfide added: Concentration of sulfide solution: 2 mg/529 ml Chlorine added:
Temperature


Reaction
Time
fmin


4.96
U r*
N
S
S IS


5.92
a
N
S
S
p

7.09
p
S
p
S
S



7.97
*
S
S

It

9.04
S tS
p
S
S


Chlorine
Reacted
(m


14.96
15.05 15.27 15.40 15.54 15.33

11.35
12.02 11.98
12.12 12.34 12,54

9.83 9,92 10.19
10.44 10.75
10.50

8.29 8.51 8.58 8,71
8.95 9.15

7.25 7.75 7.92 8.03 8.20
8.29


Chlorine to
Sulfide R+4o ( I


7.48 7.53 7.64 7.70 7.77 7.67

5.68 6.01 5.99 6,06
8.17
6.27

4.92
4.96 5.10
5.22 5.38 5.75

4.15 4.26 4.29
4,36 4.48 4.58

3.63 3.88 3.96
4.02 4.10 4.15


r --7--


525 ml.
2.00 Mg. or 3.78 ppm.
17.68 mg.
200 C.

Chlorine to Sulfide
at+I4r A


3.39
3,41 3.48
3.49 3.52
3.48

2,57
2.72
2.71 2.75
2.80
2.84

2.23
2.25 2.31 2.36
2,44 2,60

1.88 1.93
1.94 1.98 2.03 2.08

1,64 1.76
1.79 1 82 1.86
1,88


- 65 -


' - -% g-Ir oW mo &"%F,4 y


1










TABLE 13


EFFECT OF TEMPERATURE ON THE RATIO OF CHLORINE TO
SULFIDE REACTED.
SUMMARY OF OBSeRVED DATA

Source of data: Tables 9, 11 and 12.


Reaction
Time (m4, I


4.92
N
U
a
a


5.93
a




7.10
a
a
a


7.97
p




9.05
#
a
a
a It


Chlorine '?O


3.34 3.44 3.45 3.50, 3.52.
3.55,

2.57.
2.73
2.72
2.74 2.80
2.82

2.14
2.20, 2.34 2.34, 2.40 2.54

1.84, 1.87,
1.93 1.91.
1.986
2.02.

1.60 1.76 1.80
1.81 1.85 1.89-


to Sulfide Rati
(mole)
200o


3.39 3.41 3.46
3.49 3.52
3.48

2.57 2.72 2.71 2.75
2.,80 2.84

2.23
2.25 2.31 2.36 2.44 2.60

1,88 1.93
1.94 1.98 2.03 2.08

1.64 1.76
1.79 1.82 1.86
1.88


Los -Maximum 25e Devirrat on


3.38
3.41 3,46 3.46
3.47 3.50 3.50

2.58 2.70 2.71 2.76 2.80 2.83

2.20
2.18 2.32 2.37
2.46
2.54

1.87
1.90
1,92 1.90 1.99 1.97

1.72 1.81
1.83 1.87
1.91 1.93


1.5
0.9 0.3 0.9 .0.6
2.0

0.4 11.1
0,4

1 *0 .0.0
0.7

-4.1
3.2 1.3
-1.3
12.5
- 2.3

-2,2
-3.2
1.0
o4.1 .3.5


-7.3
.2.8
2.2 .3,3 .3.2
.2.6


*The pI values listed hereare the average values for the experiments conducted at the various temperatures.


- 66 -


r-1 .- 250 Deviatinn'


~"*"' "










ens 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,

tffTct of Hydrq n-qn Concentrat:kqn 9n the Reac#on

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 concentration has a very decided effect upon the reaction is clearly 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 ex-tent 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







- 88 -


allowed to progress for one minute and longer. -In this regard Particular consideration is invited to the data appear. Ing inTable 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, Here it is observed that the Intercepts or 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 vith the PHi value in such a manner that there Is a aximm slope indicated In the vicinity of PH 7.0.8 Since this slope 0Is 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 or the study would be to demonstrate the effect of the hydrogen -ion concentration on the constants a and b of the empirical equation, R = atbo 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.
1Bxpriments such as those used to determine the effect of time on the reaction were found to be most suitable for









producing the necessary data for this part of the investigation. In fact the results for the time experiments conduoted 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 all 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 A that are used in the study are summarized 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,










TABIS 14


EFFECT OF HYDROGEN-ION CONCENTRATION ON THE RATIO OF
CHLORINE TO SULPIDE REACTED OBSERVED DATA.


Volume of buffer solution:
Sulfide added: Concentration of sulfide solution 8 mg/529 ml or Chlorine added: Temperatures Ionic strength valued


525 ml.
2.00 mg. 3.78 ppm.
17.68 mg. 250 C. 0.1093


pH Reaction
Time
(min)


1
5 10 20 40, 80


5 10
20 40 80

1
5
10 20 40 80

1

10
20 40 so80


Chloride
Concentration
innm)


1200
a, 'I I
a


I
w


tI
t35
V


if S V
a
w if.i
a



1540 11
I
I
#


Chlorine
Reacted
(mr)


10.66 10.99 11.20 11.45 11,96 12.14

10.07 10.81 10.84
11.57 11.96 12.33

9.96 10.71 10.80
11.46 11.84 12.18

9.52 10.64 10.39 10.98 11.60 11.88


Chlorine to
Sulfide Ratio (ir)


5.33 5.50
5.60 5.73 5.98 6.07

5.04 5.41
5.42 5.78 5,98 6.17

4.98 5.36
5.40 5.73 5.92 6.09

4.76 5.32 5.20 5 49 5.80 5,94


Chlorine to
Sulfide
Ratio (mole)


2.42 2.49
2.54 2.60 2.71 2.75

2,28
2.45 2.46 2.62 2.71
2.80

2.26
2.43
2.45 2,60 2.68


2.16
2.41 2.36 2.49 2.63 2.69


Continued


8.35
V
V
*f S "


6,56
I,
*
*
V

if
a
a
a


6,79
-w
a
#


6.95
*

6* 9
*
V
V
a
a if
VI


\--- -~ -T- ~71 .~--r--










TABLE 14 Continued


pH Reaction
(Time
(min)


5 10
20
40 80


1
5 10
20
40 80


Chloride
Concentration
(ppnm)


Chlorine
Reacted .r(n )


Chlorine to Sulfide
Ratio (m)


Chlorine to
Sulfide
RatIp (molA


- V I / V . I I t V I. It 7i


466
a
#




213

fi ft Ifr
V
21


8.93 9.60
9.68 10.06 10.52 10.92

9*00
9 * 33 9.20 9.9 6 10.18
1039


4.57 4.80
4.84
5.03 5.26.
5.46

4.50 4.67 4.60 4.98 5.09 5.20


82.07 2.17 2.19 2.28 2.38
2.48


2.04 2.12
2.08 2.26
2.30 2.36


:|~
7.25
ft
p
r



ft
c


ft I,


" 71-


~'4'


r 1










TABLE 15


BFECT OF HYBDROGEN-ION CONCENTRATION ON THE RATIO OF
CHLORINE TO SULPIDE REACTED. CALCULATED DATA.

Source of data for calculations: able 14.
,asis for calculations - R : at where
R- Chlorine to sulfide ratio (mole),
A =Reaction time (min),
g and k are characteristic constants,


pH, Reaction
Time
(min)


1
5
10
20 40 80

1
5
10 20 40 80

1
5
10
20 40 80

1
5
10
20 40 80


Constant
A&


2.38
a
a ft ft

a

2.25
*
a ft
a ft


a
a

a ft




2.16
a
a
a
a ft


Constant


0.0333
a a a ft
W

0.0502
a
a a a

0.0473

a



0.0502
f ft

U 'ft f
a


Calculated Chlorine to
Sulfide
Ratio (molg)


2,38 2.51 2.57 2.683 2. 69
2.75

2.25
2.44 2.52 2.61
2.70
2.80

2.24
2.42 2.50 2.59 2.67 2.76

2.16
2.34 2.43
2.51 2.60
2.69


% Deviation from Observed
Ratio


-1.7
1.2 1.2 1.1
-1.1*
0.0

-1.3
-0.4
2.4
-0.4
-0.4
0.0

-0.9
-0.4

-0.4
-0.4
0.0

0,0
-2 * 9

2.9 0.8

0.0


Continued


- 72 -


6.35
a ft
a
a


64356.
a





6.79
a



89

f
a
4
ft ftEi ft. ft


Hatin tcalo










TABLE 15 Continued


pH Reaction
Time (min)


Constant
ow


Constant
k1


Calculated Chlorine to
Sulfide
Ratio (fmale


% Deviation from Observed
Ratio


1.99 0.0499 f a


2.03 0.0349
U


- 73 -


7.25





7.44

U

I Ua
p


1.99 2.16
2.24 2.31 2.40
2.48

2.03 2.14 2,19
2.25 ,.30 2.36


-3,9
-0.5
2,3
1.3
0.8 0.0

-0.5
0,9
5.1
-0.4
0.0 0.0










TABLE 16

EFFECT OF YDROGEN-ION CONCENTRA 0 o ON THE CONSTANTS OF TIE EMPIRICAL EQUATION, R: at . SUMARY OF DATA.

Source of data: . Tables Is 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.

pH Ionic Strength Chloride Constant Constant
Value Concentration a b


4,90 5.89 6,35

6.56 6.79

6.95 7.25 7.44

7.92 8.95


0.1093 0.1093 0.1093 0.1093 0.1093 0.1093 0.1093 0.1093

0.2154 0.1093


2120 603 1200

832

435

1540 466 213

0
3470


3*39

2.60 2.38 2.25

2.24 2.16 1.99 2.03 ;1.87
1.74


0.00720

0.0193 0.0333
0.0502

0.0473 0.0502

0.0499 0.0349 0.0124 0.0243


-74 -








4 75 o


FIGURE 5
EFFECT OF HYDROGEN- ION CONCENTRATION ON THE CONSTANTS OF THE EMPERICAL E9UATONJ R -ab.
CSee Table /6)


Curve I: Consfonl aCarve II : Conson/ b_


2:4


5.6 6.0


72 7p 8.0 8.4 8.8


pH


/4







- 76 -


Curve I of Figure 5 demonstrates that there is a maximum oxidation rate fbr the reaction mixture that is very sharply defined and dependent upon the hydrogen-ion concentration. This maxlutm 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 al. kaliesside,' and it appears to be destined for a value of zero at a pH somew-here 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 onehalf 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 11. 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 ionic 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













FIGURE 6

EFFECT OF HYDROGEN-/ON CONCENTRATION ON THE
RATIO OF CHLORINE TO SULFIDE REACTED

(See Tab/es 9 ~/4)


Curve I : Curve IT : Curve 12:


Reachan Recon Reac/lion Se ac//on


fimne is / min. /me is /0 min. lime is 40 min.


3634 323.0


2.8


2.4 -


2,2 -


/.8 -


50 6.0 70 .0 9
P,/










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 Ahe samej 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 11 and 111, which represent the experimental ratios obtained with reaction times of 10 and 40 minutes, respectively, It is observed that these letter curves are displaced upward in relation to Curve 1. which Is to be expected when the longer reaction times are con. sidered. 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 pff range. A significant feature of this maximum in the oxidation rate is readily recognized 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 2.6 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 8, is suggestive of the way in which the concentration of undisuociated hypochlorous 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 hypoehlorous 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 hypochlorous acid and the hypochlorite ion and the hydrogenion concentration. The method that is involved is as follows,







- 80 -


For hypochlorous acid
Hclo - I- / 0co10
a / x aC0-o = 5.6 x 10-


K 5.6 x 10-8 (29)


aHC1O 1,
18010 5 0"


a ta HC1O 010


a :a C107
aECO 10 C010 'UCI-O ' CIO-


1 : . x 0"s


10mpH
1 56 o(PH - 8)


Lt O / a1- = aT

a l/ : 1 / 5.60 x 1o(p -8) 1 0(P) - 7.282)

o-/aa 5,0 19i,(PH -8) ( H - 7.252)
S/ 5.6 x 10 - / 10PH - 7.252)
The calculated activity fractions are presented in Table 17, and the data are exhibited graphically in Figure 7. A comparison of the hypochlorous 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 hypochlorous 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


Let










TABLE 17


ACTIVITY fRACTIONS -HYPOCHLOROUS ACID

pH . aC10/ aC10"-/3.2 1.0000 0.0000
3.4 0.9999 0.0001
368 0.9998 0.0002
3, 8 0.9996 0.0004
4,0 0.9994 0.0006
4,2 0.9991 0.0009
44 0.9986 0.0014
4.6 0.9978 0.0022
4.8 0,9965 0.0035
5.0 0.9944 0,0056
5.2 0.9912 0.0088
5.4 0.9861 0.0139
5.6 0.9782 0.0218
5.8 0.9659 0.0341
6.0 0.9470 0.0530
6.2 0.9185 0.0815
6.4 0.8784 0.12386
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 0.8904
7.8 0.2208 0.7792
8.0 0.1515 0.8485
8.2 0.1013 0.8987
8,4 0.0682 0.9338
8.86 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 049989
10.4 0.0007 0.9993
10.6 0.0005 0.9995
10.8 0.0003 0.9997
11.0 0.0002 0.9998
11.2 0.0001 0.9999
14, o1,99P 9kor o


- 81 -








- 82 -


FIGURE 7 ACTIVITY FRACTIONS- HYPOCHLOROUS AC/D (See Table /7) CurvelT: Achvi/ly froclon of undlissociaoed hypoch/orous deid. Curve -I.: Acl'y Fraclion of hypoch/ore ion.





1.0





0.8











0.6





0.2




0.0I


I







do 83 -


this vicinity the activity fraction of the hypochlorous 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 S. However. in the more acid range the curves bear little resemblance to one another, The activity fraction curve of hypochiorous 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 pg.,


)jrfVe9t of _liq~t ~tnth gn the Reaction

The previous observations concerniTU the effect of
hydr ogen-ion concentration on the reaction serve to emphasize the 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 be used in most of those experiments where the use of solutions having such an tonic 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 the 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 pHff 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 be kept within the prescribed minimum of 0.05 unit. Furthermore1 it can be 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-4on concentration indicates that this small variation of 0.16 pH unit produces no Inaccuracies dut 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, and b of the empirical equation,
R = atb, were calculated for each of the experimental runs In the study, and the results are reported in Table 19. An
examination of the results indicates that the initial increase 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 Sulfide added Concentration of Chlorine added: Temperature:


solutions 525 ml
2.00 ng. sulfide solutions 2 mg/529 ml or 3.78 ppm.
17.68 mg, 950 C.


Ionic Strength


0.05463





0.07983
U U
U
U
U
0.108
V
a a a 'U
U
0.04
U


Reaction
Time
Imin


Chloride
Concentration
I( m}


7.05
U

S
U
U
U


7.02
U
p
I
p
p

6.98
U
U
*
*
p


880
p p p
U
U
* V p
1760


U p


Chlorine Reacted
(m I


9,72
9,80
10.24 10.46 10.87
11.22

10.21 10.57
11.50 11.85 12.31 12,79

10.23 10.42 11.09
11.20 11.82 11 A 22


Pi~A


Chlorine to
Sulfide In 4 tm


4.86 4.80
5.12 5.23
5.44 5.61


5,11 5.29
5,75 5.93 6.186 6.40

5.12 5.21 5.55 5.60 5.91 6.11


Chlorine to Sulfide
Rag (,mo\ )S


2,20
2.18 2.32 2.37
2.46 2.54

2.32 2.40
2.60 2.68
2.79 2.90

2.32 2.36 2.51
2.54 2,68 2.77


ppM IV, Ir ;I K ygg P


r 2 77








TABLE 18 Continued


Ionic Strength


Reaction Time


Chloride
Concentration
( m


Chlorine
Reacted I 1 1.


Chlorine to
Sulfide
Ratit (mrly


Chlorine to
Sulfide Ratio (molea


(ngn) 7pr q


0.1296
U I I
V
I
0. 1546

!1 'i

a
0.1796

a

a


I


? 00


8.95
V




8 .93
.w




6.91


a 8.89
,


a


2640
I



3520
* U 'U .1
I

4t400
a
U I I







5280
a I I I
a


10.09 10,20 10.95 11.13 11.32 11,.64
9.96
10.47 10.83 11.13 11115
11.48

10.09
10.34 10.72 10.90
10.94, 11.13


10820
10.16 10.73 10.85 11.09
11_09


5.10
5.48 5.57 5.66 5,82

4.98 5624 6.42 !6.57

5.74

5.05
6.17 6.36
6.45
5.47 5.57


5.10
5.08 5.37 5.43 5.55


2.29 2,31
2.48 2.52 2,56
2.64

2.26
2.37 2.46 2.52 2*53 2.60

2.29
2.34
2.43 2.47
2.48 2,52

2.31
2,30 2.43
2.46
2.51 2~AR







TABLE 19


E9MCT OF IONIC STRENGTH ON THE RATIO OF CHLORINEM
TO SULFIDE REACTED. CALCULATED DATA.

Source of data for calculations Table 18.
Basis for calculations: R : atb, where
'- Chlorine to sulfide ratio (mole).
t Reaction time (min),
& and b are characteristic constants.


Ionic Strength


0.05483
i U


0.07963
if
if U if if


Reaction
Time
(min)


1
5
10 20 40 80


Constant
a


2.04
if if



2.30
if
if if if if


Constant


0.0497
if if *
It if
O.0526
if
if U if
U


Calculated Chlorine to
Sulfide
Ratio (mole)


2.04 2.21
2.29 2.37
2.45 2.54


2.30
2,51
2.60 2.70 2,80
2.00


% Deviation from Observed
Ratio


-7.5
1.4
-1.3
0.0
-0.4
0.0


.0.9
4.5 0.0 0.7
0.4 0.0








TABLE 19 Continued


Ionic Strength


0.1046
U




0.1296


if


0.1548



f


Reaction
Time (mn)


Constant
A


2.23





2.28
a
a
a
U





2,27
a
a
a if
U if
2 2


Constant


0.0492
a




0.0331





0,030'7
a
a
a a


Calculated Chlorine to
Sulfide
Ratio (mole)


2423 2.42 2.50 2.59 2.68 2.77

2.28
2,41 2.47 2.52 2.58


,87 2.27 2.39
2.44 2.49 2.55 2.60


% Deviation from Observed
Ratio


-3.9
2.5
-0,4
1.9 0.0 0.0


-0,4 0.0
0.8 0.0

0.4
0.8 0, 8
,
0.8 0.0








TABLE 19 Continued


Ionio Strength


0.1798
a

a a

,0.2048
a


a
a


Reaction.
Time (min)


1
5.
10
20
40 80


5
10
20 40
80


Constant


2.28
a
a
p
V

2.29
V
p
I
a


ft
p


Constant
Sk


0.0238





0 025
a N V

a


Calculated Chlorine to
Sulfide
Ratio (mole)


2.28
2,37
-2.41
2.45
S2.49 2.53

2.29 2,39
2.43
2.47 2.52
2.56


% Deviation from Observed
Ratio


-0.4
1.3
-0.8
-0.8
'0.4
0.4

-0.9

0.4 0.4 0.0







91 -


880 parts per million, Literature previously cited (42) suggests tbo possibility that the observed effect on constant

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 with changing ionic strength, The relationships that exist between the two are represented graphically In Figure 8* It appears that

has a constant value in the vicinity of 0,050,, within the limits of experimental error, for ionic strength values 0,054630 0.0"630 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 pff value 7,92 in Curve 11 of Figure 5* All of the other points on this curve were do-









- 92 -


FI1/RE 8 EFFECT OF IONC STE GTTHH ON THE CONSTANTS OF

THE EMPIR ICAL . E ACTIONN R= a . (8ee Tob/le 9)


-4
,,)


0.070 0015 000 " 0 055 . 050 0045 0.040 0.015 0.30


0.0250,020
0


I I I It I


0. 02 004 006 008 10. 0 a.12

Ionic Sfren9#h


C)


0./4 . 46 0./8 020 022


I I








-


I !







IM 93 11


termined In solutions having a common ionic strength value of 0.1093, but this particular point In question was de-. termined at a value of 0.21544 It is seen from Figure 8 that decreasing the ionic strength from 0,21 to 0.11 has the effect or 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 pre-. supposed value,

In view or the-foregoing evidence that the effective agent In this oxidation reaction is hypoohiorous acid, it was thought advisable to investigate the effect of the concentration. or chloride ion, which was used to vary the Ionic strength values during these experiments, on the hydrolysis of chlorine to hypoebiorous acid. The equation for tis hydrolysis reaction is as follows: Cl / HCl0 / a 012/H20
Latimer (29) gives -6,315 calories per mole as the free energy change ror this reaction, which yields an equilibrium constant or 4.27 x 104, Using this value for the equilibrium constant, it Is calculated that at a pH value or 7.0 and with no chloride added all or the chlorine that was added to the working solutions Involved in the experiments existed In the solutions as hypochiorous acid. This acid was ioniz~ed in accordance with the relationships shown in Figure 7, Further calculations show that when the chloride concentration In the







.0 94 -


working solutions was 5,280 parts per million,, the largest concentration used in these experiments, the total concentration of hypodhlorous acid was decreased by only 0,22%. It is assumed, therefore, that the effects noted in this study are not due to any shift in the chlorine - hypochlorous avoid equilibrium caused by the addition of chloride ion.

@{pcq qr phlorde Copentratinp on the Rpactygn

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-







- 95 -


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 -be 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 pro. sented 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, 11 = atb, are briefly summarized for convenience as follows:
Chloride Constant Constant % Deviation of
Concentration a k _ from mean
(ppvm)
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 O. 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: 200OO mg.
Concentration of sulfide solution: 2 mg/529ml or 3.78 ppm.
Chlorine added: 17.68 mg.
Temperature: 250 C.
Ionic strength: 0.1093

Chloride Reaction pH Chlorine Chlorine to Chlorine to
Concentration. Time Reacted Sulfide Sulfide r
(pro) (min) (pm) Ratio (mg). Rato (op
0 1 6.98 9.17 4.59 2.08 �
S5 n 9.58 4 79 2.17
10 9,58 4.79 2,17
20 10.22 5.11 2,32
40 10.45 5.23 2.37
80 11.07 5,54 2.51

385 1 6.97 9.27 4.64 2.10
5 9.98 4.99 2.26
10 " 10.16 5.08 2.30
20 10.64 532 2,41
40 10.89 5.45 2147
V 80 11.80 5.90 2.67


Continued







TABLE 20 Continued


Chloride
Concentration
fit MI


Reaction
Time
(g4"I


769
a
a if
a 'I


1150
a





1540
a a



1920
a # if
a a


Chlorine
Reacted
I Il


6.97
a

a


86.95
a

if




6.98
n
a
a
6 *95 ft

it
a
a
# it ft


9,17 9.95 9,92 10.63 10.81 11.03

9,23
10.00 10.15 10.36 1120
11,84

9.52
10.64 10.39
10,98 11,60
11.88

962
10.30
10.84 11.16 11,66 12.03


Chlorine to Sulfide
Rat4 t(f


4.59 4.98.
4,96,
5.32 5,41 5,82

4.62 5.00
5.08 5.18 5.60 5.92

4.76
5.32 5.20 5.49 5.80
5 94

4,81 5,15 5,42
5.58 5.83 6.02


Chlorine to Sulfide
Rati4 (mtiA


2 .08

2.25
2.41
2,45 2.50

2*09 2.26 2.30 2,34

2454
2.68


2,16
20'41 2.36
2.49 2.63 2,69
2.18 2*33
2,46
2.53 2.64


v N , y1 qr e-Wy yq, q P


2.7


- --







TABLE 21


SPFECT OF'CHLORIDE CONCENTRATION ON THE RATIO OF CHLORINE'TO SULPIDE REACTsD. CALCULATED DATA.

Source of data for calculations Table 20, Basis for calculationsi . R atb, where
S= Chlorine to sulfide ratio (mole),
S&s Reaction time (min),
Sand b are characteristic constants.


Chloride
Concentration
(ppm)


Reaction
Time (min)


Constant
P.


2 01
p
p
V



2.07
p
P



tp
*
U


385
rr

if V U


Constant
k


0.0477

I


O. 0523
a U p p

I'

I


Calculated
Chlorine to
Sulfide
Ratio (mole)


2.01 2.17

2.32 2.40
2.48

2.07 2.25
2.34 2.42.
2.51 2.61.


% Deviation from Observed
Ratio


-3.4
0.0 3,2 0.0
1.3
-1 2

-1.4
0.4 1.7
004 1,6
-2.3


Continued


1 '




Full Text

PAGE 1

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

PAGE 2

TABLE OP CONTBN'l'S INTRODUCTION AND mwmv OF THE LI'J.'ERATURE Page 1 EXPBRIMBH!AL METHODS . •. 14 Preparation and Standardisation ot Solutions ......... 14 Description ot Appara tua . ........... 23 ; , . . Experimental Proced~ea 30 . . BXPBRIMBH'J'AL DATA . •. •. •. •.. 42 . . . . . t>recision or Determinations 44 ; Bf"f"eot ot Concentration on the 47 Etteot of" Time on the 55 Bf"t"ect or Temperature 9n the 59 Et.feet or Hydrogen-ion Concentration on the Reaction 67 . Bt"f"ect ot Ionic Strength on: the Reaction 83 Ef't"ect or Chloride Concentration on the Reaction 94 DISCUSSION OF RESULTS 101 BIBLIOGRAPHY . .. 111 ACKNOVLEDGEME.NTS •. •~. ~. 115 BIOGRAPHICAL tUHS .. .... 116 COMMITTBB RBPORT ........ ~. .... 117

PAGE 3

INTRODUCTION AND REVIEW OF TBE LITERA 'lURE Sulrides are commonly round in the natural waters of several regions or our land, where their occurrence is noted with tar greater frequency in ground waters,, which are generally devoid or oxygen, than in,aurrace waters. The state or Florida lies in a region where sulfide concentrations in natural waters up to several parts per million are not un common. The origin, or these sult'"ides is attributed to bio-. logical and chemical processes whereby sultur eompounds are decomposed and sulf'ur in a tree or oxidized state is reduced. Experience has: taught us that the presence of sulfides . in a water renders that water highly undesirable tor domestic and industrial usages. 'lhere have been numerous instances where the presence or sulfides in an industrial or domestic water supply hAs been round to be responsible ror orrensive odors, exceasive corrosion and conditions leading to heavy growths of' mioro-organisma-vith the undesirable consequences attached thereto. Furthermore, sultidea have been round to be responsible tor extensive damage to greensand zeolite beds. The result or such observations has been the realisation of" the urgent need tor theremoval of' su1.r1des trom waters to be used ror domestic and industrial purposes. It has been common practice f'or many years to ef't'"eot the removal or sulfides rrom water supplies by aeration. By this 1

PAGE 4

method part ot the sulfide, the amount being largely depend ent upon the hydrogen-ion concentration ot the water, is ex pelled to the atmosphere as hydrogen sult"ide, and the remain der is oxidised to some degree by the disaolYed oxygen in the aerated water. The method is attended by many ditticulties ot: an engineering nature~ and its etticiency is dependent up on seTeral chemical and physical Yariables. Theretore. the over-all ettectiveness ot the method may be seen to vary con siderably trom one installation to another. Chlorination as a means tor sulfide removal trom natural waters is or rather recent origin. It appears to date trom 1926~ (1) in which year it vas resorted to at Be-Yerly Hills, California, atter aeration alone had been proven to be in ett"ective tor complete sulfide remoTal. As late as 1936 Cox, (2) in reviewing progress in the elimination ot ta&tes and odors trom water supplies, expressed interest in the apparent novelty of hydrogen sulfide removal by chlorination when he made the :following statement: •Mention should be made, hove•er, to the practice at Holland, New York, where aeration is used to remove hydrogen sulfide from well water. The Bolland supply is also disinfected with chlorine, and it 1a interesting to note that the chlorine will react with any residual hydrogen sultlde present in the aerated water and completely remoTe it.•. Earlier mention is made o:r the e:r:rect ot hypoohlorite on hydrogen sultide in septic sewage liquors. Rideal, (3) in reporting the results at a long

PAGE 5

series of' experiments dealin,; with hypochlorite treatment ot septic sewage liquors in the 1908 report of' the Royal Com mission on sewage# calls attention to the tact that to 5~ treatment vith hypochlorite solution is sutf'icient to do ava7 with theof'f'ensive hydrogen sulf'ide smell leaving in its pl~ee a comparatively inottenaive odor.of' spent bleach and f'resh sewage•. Discrepancies in the literature in regard to the dosage of chlorine required tor sulfide removal indicate the need f'or tundamental research concerning the reaction between sulfides and chlorine in dilute aqueous solutions. In the usual case in the literature it is assumed that chlorine oxidhes the sulride to free sulfur, and the dosages required are based on the stoi.chiometrie relationallips pertaining to this reaction. Hoover (4} states that the reaction between chlorine and hydrogen sultide is as rollovs: H 2 S ,' Cl 2 : 2BC1 ,' s. He concludes that tor the removal or each pound ot hydrogen nlfide 2.1 pounds ot chlorine are required. Pomeroy and Bowlus (5) ha-ve the :following to aa7t •When ,. . .,, . ' chlorine is added to a pure vater solution or sulfides, .it ,, . ' . re~uires only 2.22 parts or chlorine to destroy one part ot sulfide as indicated b:, the rolloving reaction: c1 2 ,' u 2 s: S ,tHc1.•. Thus it is seen that these authors are basing their dosages on the stoichiometric relationships involved in the reactions they ha.Ye written. However. Pomeroy and Bowlus (5) in determining the chlorine dosages just sufficient

PAGE 6

-4to eliminate the sulfides f"rom Yarious samples or septic sewage round that the ratio or chlorine to sulf'ides varied from 3 to 9, aTeraging 5.3. They explain the higher dosages in the case of sewage by pointing out that other reducing agents would consume a part or the chlorine. Experiments by Powell and Ton Lossberg (6) indicate that the chlorine dosage for removal of hydrogen sulfide. trom some natural vaters in which they were interested approximates the stoichiometric . value dictated by the folloWing reaction: 401 2 /. 4H 2 o /. u 2 s: u 2 so 4 /. 8BC1. The stoichiometric value calculated from the relationahips involved in this reaction is seen to be 8.84 parts or chlorine to one part or hydrogen sulfide expressed as sulfide. In actual practice chlorine dosages for sulfide remoT&l are generally determined by chlorine . demand teats (7) with the particular water in question. How ever, Wallace and Tiernan Co., Inc., has patented a method (8) whereb7 the amount or chlorine to be added is indicated by means or a ehange in an electrode potential as chlorine is added to the aqueous material. These methods require no knowledge or the reactions involved in the oxidation oft.he sulfides. The oxidation ot aultides 1n aqueous solutions has been studied by many investigators using a variety or oxidieing agents. A review or the work or several or these investi gators indicates that the products or the oxidation •ary with the relative oxidizing potentials or the oxidizing

PAGE 7

5 couples used. In the tollo1d.ng discussion couples with oxidation potentials in the vicinityot or greater than one wlt are considered to be .strong oxidants, and those with potentials less than one volt are eonaidered to be weak oxidants. Among the weakest oxidants studied were neutral to alkaline solutions or suspensions or potassium chromate (9,10), potassium dichromate (11)" ammonium chromate (12), lead chromate (13), silver chromate (13),, merct1rous chromate. (14), barium chromate (15) and solutions or various water soluble aromatic nitro compounds (16) •. The principal pro ducts or the oxidation or alkali sulfides and hydrogen sul fide.by these oxidants were round tobe polysulfides,, tree sulfur, thiosultates, sulrites and sul:fates. The f'ormation or sulfites and sul:fates appears to be dependent upon the temperature and alkalinity of the solutions, higher temper atures and lover hydroxyl-ion concentrations tavoring the production ot sulfates. Bullock and Forbes (16) show that the oxidation ot &Ultides by such mild oxidants as aromatic nitro compounds progresses only as tar as t"ree sult'ur. They attribute the presence ot the other final products, thiosul tatea and polyaultidea, to a secondary reaotion between •active• sulfur and hydroxyl-ion, which they write as tollowa: GOH/. 128 s 8 0 3 --.; 28 5 ~/. 3H 2 o. The investi• gators-describe •active• sulfur as that sulf"'ur set :rree by an oxidation or a aulf'ide or in some similar wa:,. This work of' l3ullockand Forbes brings to,mind the oxidation or hydro

PAGE 8

;.;, 6 genaultide by iodine,. which is considered te be a weak oxidant. That oxidation proceeds quantitatively to f'ree suU'ur in an acid medium (17) and is the basis ror a well known method tor the quantitative determination or hJdrogen sulfide in aqueous solutions. However, in alkaline solutions nltate appears as a product or the oxidation (17, 18). Some JDGderately weak oxidants that have been used in studies or the oxidation or 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 tr1ox14e (24). The principal products resulting from the oxidations with these oxidants are free sulfur, thiosultates, aulfitea and sul fates. It is noted that in these stuties polyaultides are not mentioned among the products found. Strong oxidants that have been employed include oxygen in aeid solutions (20)~ potaasium iodate 1n acid solutions (25), nitric acid (26), paraper1odic acid (27) and a neutral solution of potassium permanganate (28). Pree sulrur and sulfates were found to be the only end-products With these oxidant• except in the case of' the oxidation with neutral potassium permanganate~ in which case it was f'ound that under certain conditions relative to concentration• a dithionate occurs as a f'inal product. The preceding reYiev or t.he work of' some of' the investl

PAGE 9

7 gators who have studied the oxidation or sultidea in aqueouu solutions inclioatea that the hydrogen-ion concentration or the medium in which the oxidation takes place is a very important raetor 1n determining the characteriatios or the reaction. Free aultur is app_arently the primary oxidation product or sulfides and may be the only end-product 1n aeid solutions when weak oxidants are employed (16,. 17}. In basic aolutions there are obanges noted in the tinal products ot the oxidation (17, 18, 20). These changes are evidently ex plained by ef'f'eots ot the hyclrogen-ion concentration on the. oxidant, on the eultide equilibria relationships and on the . ' " . intermediate products of the oxidation, The oxidation potentials ot the large majority or the oxidants used by theee investigators are dependent upon the hydrogen-ion concentration ot the medium (29). Por example 1 it is noted that the standard oxidat1o~ potential hr the water-oxygen couple in acid solutions is -1.229 •olts. vhich lncli.catea that oxygen is a strong oxidant in acid media. The potential f'or the same couple in a neutral solution 1a only -0.815 'Volt, indicating a moderate3:y weak oxidant. Another possible etteet ot the hydrogen-ton concentration on certain oxidants is illustrated in the apecific case where nitrous acid was employed as the oxidising agent. Mention was :made by Bagster (21) that the results ot the study suggest that tree nitrous acid is the active agent rather than nitrite ion. If' this is true the importance of'

PAGE 10

8 . the hydrogen-ion conoentration lies in its ini-luenee on the equilibrium involving tree nitrous acid and nitrite ion in the reaction mixture. Still another erteot or the hydrogenion concentration on an oxidant which leads to subsequent changes in oxidizing ebaraeteristics is: noted in the ease or . iodine (17). Although the oxidation potential or the iodide iodine oouple :ls not dependent upon the pH or the solution at pH Talues below 8, at higher pH values the iodine Will react with the hydroxyl-ion to yield hypoiodite (30) thereby alter ing the oxidizing properties or the mixture. It is interred rrom the work or Kapp (31)_ vho deduced rrom experi.Jnents dealing Vith the oxidation of alk-.11 ~1r1c1es that hydrogen sulfide 8hould be more easily oxidiaed in aqueous solutions than sodium sulfide, that there might be a possible ettect or hydrogen-ion concentration on the oxidation of' sulfides trom the standpoint of' its influence on the aulride . equilibria relationships. The influence or the hydrogen-ton concen tration on intermediate oxidation products or aulfides maybe exemplified by the reaetions between sulfur and alkali and alkaline earth hydroxides in aqueous solutions. Tartar (32) round that the primary reaction is as follows: 60H-J 8S = 2s 3 -/. s 2 o 3 /. 3820. An excess or sultur was found to yield pentasulf"ide by a secondary reaction. A• pointed out by . Bullock and Forbes (16) the reactions do not proceed rapiclly at 25 with ordinary rhombio aultur . , but when the aultur ta in the •active• state the reactions are quite rapid

PAGE 11

-.9 . even at th:J.s temperature. A turther example or the eftect or h7drogen-1on concentration on a product or the oxidation or sultides is seen in the stability exhibited by thiosul rates 1n alkaline solutions (33, 34). Now let us consider some or the properties or chlorine as an oxidizing agent in aqueous solutions. Latimer (29) gives -1.3583 volts as the best value tor the standard oxidation potential or the chloride-chlorine couple, but he points out that the potential becomes meaningless in alka line solutions because or the hydrolysis or ~hlorine and the formation or hypochlorite. He lists the standard oxidation potential or the chloride-hypochlorite couple as ~0.94 volt. From this latter value it is calculated that the approximate values or the oxidation potential or chlorine in aqueous solutions having pB values or 9 and 7 are -1.24 and -1.35 volts., reapectivelJ. Th11s it is seen that chlorine is a strong oxidant throughout the normal pH range or natural waters. Higgins (35 .36), Rideal and Evans (37) and Reming ton and Trimble (38) have studied the ettect or acids on the oxidizing properties or hypochlorite solutions., It is apparent :trom the works or these investigators that the oxidizing power. or hJpochlor1 te solutions ean be markedly increased by the addition or weak acids. JU.deal (39) sug gests that tree hypochlorous acid is the ertective oxidizer in such solutions and attributes the ertect or the added acids to the resulting increase in concentration or that com

PAGE 12

10 ponent ot the ox1di9ing mixture. Rideal and Evans (37) call attention to the Talues or the dissociation constants ot hypochlorous and carbonic acids to show that tree oarbonio acid will liberate tree hypochlorous aeid trom hy.pochlorite solutions. Higgins (36) points out,that whereas the addition or an excess ot boric acid to a hypochlorite solution results in a solution ot very energetic bleaching properties, the addition .or an excess or hydrochloric acid gives a solution ot very weak: bleaching properties. He also attributes the oxidizing properties or such solutions to the active mass or tree hypoehlorous acid and explains that the. boric acid liberates tree hypoohlorous acid while the hydrochloric acid liberates tree chlorine. Weiss (40) is discussing the kinetics or chlorine bleaching claims that the active agent 1• chlorine monoxide or undissociated hypoohlorous acid. He mentions ano~her interesting point in connection th the oxidising activities ot chlorine solutions# and that is that there is an apparent maximal rate of" attack on cellulose :fibers at a certain pH value, the rate talling rapidly rith either an increase or decrease in pH. Blakely (41) measured the oxidation potentials or hypoehlorite solutions having pH values f'rom 2 to 13 and.round that a maximum is indicated at pH 7.o. It was discovered by Higgins (42) that chlorides have an aceelerati;;; ett-ect en the bleaching action o:r chlo rine solutions. The ettect was round to be an :lmmediate one, atter which action the solutions behave as though the chlo

PAGE 13

-11 rides were not present. Chlorides produced by the reduction of the hypochlorites during the bleaching reaction appear to have a negligible accelerating effect. It is mentioned by Higgins (43) that there is a secondary reaction between hypochlorous acid and neutral chloride whereby nascent chlorine or energetic bleaching properties is produced. " Information in the literature in regard to the oxidation or sulfide solutions by chlorine appears to be rather meager. Stock (44) reports that the oxidation or a dilute hydrogen sultide solution 1n anhydrous, liquid hydrogen chloride yields sulrur as the oxidation product, which lends support to the view that 8Ulf\u' is the primary product in the oxidation or such sulfides. Perel'man and Lelyakina (45) found that hydrogen aultide in acetylene gas is quantitatively oxidized to sultate when passed through a solution ot sodium hypoehlo rite and propose a quantitative method for the determination of' hydrogen sulfide in acetylene based on this reaction. Some comprehensive work has been done by Choppin and Faulkenberry (46) on the oxidation of' aqueous sulfide solutions by hypo chlorites. 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 tact that the end-products or the oxi4ation are ault'ur and sulfate* the ratio depending upon aucb factors as relative concentrations of' the original reactants, hydrogen-ion concentration of' the reaction medium,

PAGE 14

12 temperature, standing time and rate or addition ot reactants. The stand is taken that aulf"ur 1s the primary product or the oxidation, whereas sulfate results as the end-product or secondary reactions that may occur simultaneously with the primary reaction. The errect ot the above-mentioned tactors on the ratio or sulfate to sulfur produced is attributed to the influence or these raetorson the secondary reactions. For example, higher temperatures were round to increase the proportion or sulfate, and pH values or 13.8 or higher were round to result in a quantitative oxidation to sulfate. The result in each or these cases is explained by the investi ~tora on the basis or the equation, 60B/. 8S: as 3 /. s 2 o 3 /. 3H 2 o, Where increased temperatures and high alkalinities ~avor the solution of sulf'ur (32, 16). The pro ducts are subsequently oxidized to sulf'ates (47). It was also noted that there is a quantitative oxidation or sulfide tc sulfate at pH values of 2 or leas. The explanation given tor this depends upon the presence or chlorine monoxide in acid solutions or hypochlorites (48) and its t'unction as a reagent tor. the re-solution of colloidal sultur. Between pB values 2 and 13.8 the proportion of sulfate was found to de crease to a minimum at a pH value in the vicinity of 10.

PAGE 15

ST.A!fffflBNT OF PROBLEM . The object of this investigation vas to make a study or the oxidation or sulfides in very dilute aqueous solutions, sueh as those normally encountered in water works practice. Particular attention was devoted to the ettect on the reaction or hydrogen-ion concentration, time, concentrations or react ants, temperature~ ionie strength and chloride concentration. Since the study was made with a view toward application in the water treatment field, the limits adopted tor the various variables have corresponded where posaible to those commonly experienced in the water works field. 13

PAGE 16

EXPERIMENTAL METHODS The methods chosen tor use in these investigations were based upon the familiar chlorine demand method (7). Sulfide solutions or various known characteristics vere accurately made up 1n the absence of oxygen in a series or reaction vessels, known dosages or chlorine vere added, and the ' amount or chlorine or sulfide remaining after a measured elapse or time under controlled conditions was accurately determined by iodometrie or iodimetric methods, respectively. Prom the results or the residual ehl.orine or sulf'ide determinations it is possible to calculate the values or the ratio or chlorine reacted to sulfide reacted. The details or the procedures employed are given later. I f.!:ftpar1ti9n and St!:Jldatd&Hiion of SoJutisms Stand.ant .Ll.l. pqtpssiP!!! die!ifopt:e !tJuttPll• Some reagent grade potassium dichromate was pulverized in an agate m~rtar and dried in an oven at 150 8 . 200 c. 9.808 grams ot the material were dissolved in distilled water and diluted exactly to 2 liters. ;tandftrd 0 1 {)1 I pptas1ium d;!cbro9te aolut&on. 200 ml. or the standard 0.1 N solution were transferred to a 2 liter volumetric tlask by means or a calibrated 100 ml. volumetric p1pet, and then the solution vas diluted exactly to 2 liters. 14

PAGE 17

15 SuJ.[pric A2!!• Concentrated reagent grade sulturic acid was used where this acid was called tor in the pro cedures. Poia11&um iodide. Reagent grade potassium iodide crystals that have been tested tor the absence of iodate" were employed in the investigation. S;&eroh ind'7ptpr 12J.ution. 5 grams or potato starch were mixed with a little cold water in a mortar and ground to a ~bin paste. The mixture vaa poured into a liter or boiling distilled water, stirred and allowed to settle over night. 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. Zin9 aeetai, solution. 240 grams or reagent grade zinc acetate were dissolved in one liter or distilled water. 0,01 I. iodine. soluti2n. 2~54 grams or reagent grade iodine were dissolved in several ml .. ot water containing 8 grams or ioclate-f'ree potassium iodide. This solution was diluted to approximately 2 liters. No standardization was necessary. Hfdtoohlortc 1c1d. Concentrated reagent grade hydro chloric acid was used where this acid was called tor in the procedure. Acetic acid. 500 ml. or glacial acetic aeid was diluted to one liter with distilled water.

PAGE 18

16 J!.al. l! sqd;&g :&bJ.S!•Dltatp sglut1on. Appro:x:ima tel1 200 grams of c. P. sodium thiosulfate pentah1d.rate were dissolv ed 1n 8 liters of recently boiled# cooled distilled water containing o.s gram ot sodium carbonate. A f'ew ml. of chloroform were added as a preservative, and the solution was allowed to stand for several days before standardization. The solution vas standardized periodically against standard 0.1 N potassium dichromate in the folloWing manner. To 300 ml. of' distilled water vas added, with constant stirring, 2.5 ml. of' sulturic acid, 25.oo ml. of' the standa:rd dichromate solution and 2 grams ot potassium iodide. The mixture was allowed to stand f'or 6 minutes in dif'f'used light and then . titrated with the thiosulf'ate solution, starch being used as the indicator .!!&.2l, I. sodium thiosultate golution. This reagent was prepared by diluting a measured amount of the aged and standardized 0.1 N sodium thiosulf'ato solution with freshly boiled and cooled distilled water. A few ml. of' chlorof'orm were added as a preservative. The solution was standardized datlY in the following manner. To 300 ml. of' distilled water was added, w1.th constant stirring., 2.5 ml. of' sulf'uric acid., 25.00 ml. of the standard.O~Ol N potassium diehromate solution and 0.5 gram of' potassium iodide. The mixture was allowed to stand f'or 6 minutes in diffused light, and then titrated with the thiosulfate, starch being used as the in dicator.

PAGE 19

17 . Chlorine va;tff• Chlor~ne gas vas alovly bubbled, with ocoaaional shaking. through . s .. 9 liters of' distilled water contained 1n a black enameled bottle fitted vith a siphon. Samples of' the water were taken every rev minutes~ the chlorine concentration determined in accordance vith the s~andardization procedure . , . and . bubbling was discontinued when the chlorine concentration reached 1.1 1.2 milligrams per milliliter. When . the atmosphere . vas excluded f'"rom the . . . . . sol~tion by a aimple . check : valve the solution was f:'ound . io retain its strength f'ordays., losing strength at a rather ; . ') . constant rate of' approximately : 0.016 mi~ligrams per milli liter per day. Standardisation . was accomplished daj.lf 1n the following manner. To . 275 ml~ of' distilled water was ad,ded with •~rling., 10 ml. of' acetic acid and 0.'15 gram of' potassium iodide. 50.00 ml. of" the chlorine water was measured beneath the surface or the solution., and it was titrated imme~iatel:, vith standardized . 0.1 Nsodium thioaulf'ate vith starch as the indicator. Stqck sulfide solu\&on. . 2 11 ters of' distilled water ' were boiled at a moderate rate f'or 20 minutes in a 2 liter ' . Erenmeyer f'la.sk and then cooled in a vater bath under a . . . . ' nitrogen atmos~ru,re. 7.1 grams o . f' reagent grade Na 2 SH 2 o crystals, f'"reshly washed and . blotted dry., , were dissolved , in the water, care being taken to prevent the en , trance ot , air into the flask. The resulting solution had a sul:fide concentration of' approximately o.5 milligram per milliliter,

PAGE 20

18 and the concentration waa found to remain f"airl7 constant as long.as tlie soluiion was stored under an atmt)aphere ot nitrogen~ The solution was standardiz~ci da11, :1.n accordance . . . . . . with the t'olloWing procedure. . 10.00 ml. of' the sulfide aolution were measured f'rom a mlcro-buret beneath / the surtace . _, ::: or 25 ml. or zinc . acetate solution) and the . resuliing mixture . . .. was diluted vi th 110 ml. or distilled water. 00.00 ml, or 0.01 N iodine solution were pipeited in, . the , soiution vas . ' / ' I \ acidified with 5 ml. or hydrochloric acid, and then it was , ' \ allowed to stand tor 6 minutes fit diffused light. The mix~ ture was titrated with standardized 0.01 N ' sod:lum . thioaulf'ate# .•. using starch as ' the indicator. A blank determination was carried out . on the reagent•~ Dj.luting water. The diluting water was prepared separately for . each ' run by 0 boiling 4.5 liters or ,distiiied . . . ' . ' . .. wa ' ter . for 45 minutes in a 5 ter;, round~bottom 1 boiling f'lask and cooling in~ water bath under an atmosphere ot nitrogen . A water resulted that was practically free f'rom oxygen, and ' it . was stored under a pressure of nitrogen. : S _ t~ndartrbJfet 10Iutio~•~ Clark and Lube but:-c,r solutions haTing pH values ot 5.,00, 7.,00 and 9.60 were . care• fully prepared in accordance with the . instructions given by Clark (49). These solutions were used . in the invesitgation as standards by means of which the pH meter was calibrated at frequent intervals. Concentrated butf'er solutions. These buf'f"ers had to be

PAGE 21

19 so constituted that the addition of a measured volume to a definite volume ot diluting water would give a working buf"f'er solution or certain known characteristics, and consequently their preparation presented many problems. Let us :first -mention the characteristics desired in the working butter solutions, which may be defined as those solutions that were prepared in the reaction vessels prior to the addition of' the reactants to tix the conditions under which the reaction was to progress. It was thought that tor convenience these. solutions should have the same volume in every case through;.;,; out the investigation. The most convenient volume to work with was round. to be 525 milliliters. Furthermore, it was considered that except in the ease where the errect or 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 or buf"ter mater ials 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 round that the most convenient combination of' volumes to use involved the addition of' 25 ml .'of' concen trated buf"ter solution to 500 ml. of' diluting water. There fore, the concentrations or the concentrated buf"f'ers were based on this combination of' volumes where ever possible. In the case or the pH 5 butter solution the limited solubility or the butter materials interf'ered-v:l.th this plan. and a con

PAGE 22

20 centrated butter of one-halt the calculated strength was employed. This means that it was necessary to add 50 ml. of' this. particular concentrated butter solution to 475 ml. or diluting water to give the previously decided upon 525 ml. or vorkingbutfer solution. I ,The materials used in the preparation or the concen trated buffers were those or the Cl.ark andLubs series (49),. with potassium, chloride being added in addition to equalize ' ' , the ionic strength values or the working butters. Table l gives the compositions of_the various concentrated butter solutions employed during the course of'the investigation as well as the ionic strength values and chloride concen trations or the working butf'er solutions derived from these c.oncentrated solutions. The solutions were made up in nitrogen-rilled volumetric flasks trith 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 1n .the form or a carefully standardized, oxygen-rree solution that had been ~tored under a pressure ot nitrogen. The finished butter solution was tranaf'erred to a nitrogen-tilled storage bottle and was kept under a nitrogen ntmosphere. Some apprehension was felt over the possibility that the phthalate salts in the pH 5 buffer might be chlorinated during the runs inYolving this particular buffer. HoweYer,.

PAGE 23

pH or working butter . a9lut:l.on .. 5.0 6.0 6.4 6.6 a.a 7.0 . 7.2 7.4 s.o . 9.0 I refers TABLE 1 COMPOSITION OP CONCENTRATED ~OFFER SOLUTIONS AND CHARACTERISTICS OP WORltllfG Bl1P.FER SOLUTIONS~ Components or oonc•ntrated,butterMl/525 ml Charaoteristioa ot sol\ltiona (gram1L11terl working vocktng bua:@t solu;&ipn1 IOJ.C 8 Ht0 1 KH&POi n 3 Boa NaOH KCl buffer •I witb.out Added 01I Witt Q 1 added KCJ. (ppm). add@dl'.01 107.2 -----........ 10.02., .59 50 o.04885 2140. 0.:.1093 '. ' .......... 214.4 ----7.i84 20.,s,7 25 0.,09210 608 0.1093 -----142.9 ... --10.68• 53.33 25. . 0.01520 1208 0~1093 ------142.9 ----~ 14". 95 37,..,05. .25 0.08560840 o •. 1osa .. .. _____ 142.9 ......... ,. 19.70 19~36 25 0.,09690,, 438 0.,1093 ......... 71.45 ---12.44 ........ 25 0.05463 ---o~.0546:; ....... 114.4 ___ ,..,;.; 23.52. 20.77 .25 o.09aoo 470 0.1093 -----114.4 --26.54 9.492 25 0.1032 215 0.,1093 _ ........ 214.4 ---58.96 . ....... 25 0.2154. . ....... 0.2154 ........ ........ 32.47 s.948 154.4 25 0.,03585 3500 0.1093 , to.ionic strength value. I rd ""'

PAGE 24

"'"22 . . the cnlorine losses in the blank determina tion _ s vi th this buffer solution were of' the same order of magnitude as those with the phosphate and boric acid solutions. Therefore., it is concluded that no error due to such a chlorination has been introduced. Pgtassiwn Chloride solptions. Potassimn chloride solutions were employed in those parts of' the investigation . . dealing vi th the et"t'ects of ionic strength and chloride conoen tra tion on the oxidation. . They were used as a Jlleana of "Vary:• 1ng these two -variables. Por the ionic strength experiments a potassium chloride solution vas desired of such strength . that each 10 ml. substitution ror diluting water 1n the 525 ml. Yolume of working buffer solution would increase the ._ ionic strength value of the working solution by 0.0 _ 25 units~ Such a solution vas calculated to contain 97.86 grams of' potassium chloride per liter. In the case or the chloride experiments a potassium chloride solution was wanted or such concentration that 5 ml. substitutions for concentrated butter solution (pH 7.0) in a series or working buf~er solutions originally containing 50 ml. or eonoen trated butter per 525 ml. or solution would yield a series or working butfer solutions having a constant ionic . . strength -value of' 0~1093 and chloride concentrations varying by equal increments trom zero to approximately 2,,000 parts per million. Such a solution vas calculated _to contain 85.52 grams of' potassium chloride per lJ.ter.

PAGE 25

.. 23 These potassium chloride solutions were prepared in nitrogen-filled volumetric flasks with a great deal or attention being devoted to the exclusion or .air. The di.still ed water used in the solutions was rendered oxygen-:free by boiling tor 30-40 minutes and cooling under nitrogen pressure. The finished solutions were transferred to ni.trogen-tilled storage bottles and were kept under an atmosphere or nitrogen. Descript!gg pt A1pamtu1 The reaction vessels used in this investigation were con structed trom 625 ml., wide-mouth, amber glass, reagent bottles such as those commonly used in the packaging or laboratory chemicals. The lids are or the screw-cap variety and were molded f'rom a plastic material. Conversion of the bottles to reaction vessels for the experiments involved the drilling or two holes, 6.5 mm. and 9.5 mm. in diameter, in each of the lids and the cutting of a gasket tor each lid f"rom a sheet or rubber packing about 2 mm. in thickness. The holes were drilled about 1 cm. f'rom opposite edges of the lid and on a line passing through the center. They were titted with rubber stoppers that had been earef"ul.17 ground to the proper size to insure a positive closure. The gaskets were out in the shape of' a doughnut to tit 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. Vhen the lids of' the vessels were screwed down firmly against the gaskets

PAGE 26

24 and the holes in the lids were stoppered the vessels were gas tight. Figure 1 is a schematic diagram showing_th~ nitroge~ assembly that was used to till the reaction vessels with nitrogen, to prevent the entrance or air into the vessels dur ing the preparation .of' the working solutions andto.maintain an atmosphere or nitrogen over the various stock solutions. The source or nitrogen was a steel cylinder (A) containing.com pressed nitrogen that was originally under a pressure or 2,200 pounds. The f'low or gas from this cylinder .was regulated by-, . needle valve (B) whi'm. was equipped with a pre~sure gauge~ From the needle valve the nitrogen flowed through rubber con nected tubing to a pressure regulator (C), to solution storage bottles (D) and to a manifold (E). The pressure . regulator consisted of' a 100 ml., tmgraduated, glass cylinder that contained about two. inch,s or mercury,. and was stoppered by a 2-hole rubber stopper through which passed the nitrogen tube and the leg or a trap. The nitrogen tube extended to the bottom or the eyli~der. This arrangement provided tor a constant pressure or nitrogen in the system and allowed wasted nitrogen to escape without danger or air diffusing into the system. The storage bottles (D) were a series or 2-liter. wide-mouth bottles interconnected by rubber and glass_. tubing and containing concentrated butter solutions, potassium chloride solutions and sodium hydroxide solut:ton.. They were: provided with rubber stoppers through which stoppered, glass

PAGE 27

--FIGURE I SCHEMATIC DIAGRAM OF NITROGEN ASSEMBLY -1!-
PAGE 28

26 sleeves protruded. These sleeves were tor. the admittance, of pipets so, that •the solutions C1.luld be withdrawn whil_e a •~ream of. ,nitrogen kept back the air. The manif'old (El was cone:tructed from 2.5 cm. tubing abo_ut five feet in length, and it_ .vas ~itted with nine, equally spaced:, 0.5 cm. nipples to vhieh rubber connections.could be made •. Two or these ' . ' ' ' nipples served to connect nitrogen pressur_e_ . to . the diluting water _and. _stock sulf"ide. solution storage vessels •. Connected to . th_e_ remaining seven nipples. were. glass ant'.{ rubber ,tubes that were u,ec1 f"or tilling the reaction vessels, (Pl with ni t~ogen. . :During the filling process . the vessels. were: . supported in an inverted .position on n wooden raek.of".simple design (G) whic:tl held .seven or. the jars. A 1 liter .. beaker ' ' . . {H) was. used to. catch .the water f"oroed rrom. the reaction . vessel by the.nitrogen during the filling process. ~own in-Figure 2_ is a_schematie dingr,am.of' the apparatus that was used in the preparation of' the working solutions in the nitrogen f"illed reaction vessels. The diluting water was stor.ed iJl a 5 liter.,. round bottom. boiling flask (A) which was provided with a 4 hole rubber stopper. Through the to~ hol.es 1n this stopper passed a nitrogen tube .trom the manifold (E), a_ rubber stoppered sleeve that served as. a vent, .. a siphon to a.• volume'tric pipet (B). and a return tube f'rom the top of' the pipet. This return tube closed the system and prevented exposure of' the diluting water to air. The pipet (B) vas a 500 ml., semi-automatic, volumetric pipet eapeoiall;y designed

PAGE 29

27 .. FIGURE 2. SCHA1ATIC DIAGRAM OF APPARATUS I/SEO PREf?ARATIOlv OF WORKING SOLUTIONS. IN G ! ----i i ~===:t:::=== ---'-'-----'~-----' ~-----.J;.--"---~ ---....il---'-L--'-'--'---L---~-~_u_ 0 Legend A Diluf ing waler sloroge /-on/: B 500 ml. J semioofomol/c} Y0/1/mefr/c pipe/ C Srock sv/ride so/vi/on sforoge f'los); D /() ml micro buref f Nilro_gen 1nonil'old r -Ret7clhn Yes.sel G -Chlorine waler slor<7ge bollle H fl -50 ml. bore/ I /0 ml rYJ/cro bvref I I

PAGE 30

28 and constructed ror this investigation, and it was calibrated to deliver 500 ml. at 25 degrees. Provision was made in the arrangement or the apparatus so that during the measuring or the diluting water into the reaction vessel (F) and during the subsequent additions of solutions a tubewas available from the nitrogen manifold (E) to deliver a stream of nitrogen into the vessel and thus keep back the air. The stock sulfide solution was kept in a 2-liter, erlenmeyer f'lask'(c) that was fitted with a 3•hole rubber stopper. A nitrogen tuberrom the manifold (E), a stoppered vent tube and a siphon to buret (D) passed through.the.three holes. The buret (D) was a 10-ml.,. micro buret to which a side arm had been added just above the stopcock. A nitrogen filled, rubber bladder of the type used in an Orsat-gas analyser closed the top of the buret and pre vented the exposure of the sulfide solution to air. A 10 liter, black, enameled bottle (G} with an outlet near its bottom eerved aa storage tor the chlorine water. The bottle was stoppered vith al hole rubber stopper through which pro truded a simple check valve that acted as a barrier to the diffusion of gases to and from the atmosphere. The cheek valve consisted of a short length of glass tubing stoppered at one end with a 1-hole rubber stopper through which was inserted a short section of smaller diameter glass tubing that had a robber policeman stretched over its lower end. A slit in the policeman allowed air to enter the bottle only . ' when solution was being removed. The ehlorine water storage

PAGE 31

29 bottle was connected to a 60 ml. buret (H) and a 10 ml. buret (I) by gum rubber connected glass tubing. The 50 ml. buret was or the semi-automatic type that is operated by a 2-vay 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 elamp. The thermostat used in this investigation was ot the conventional, manually operated type. The water bath con sisted or a box 20 inches long, 13 inches lfide and 11.5 inches deep that vas constructed ot o.75 inch, cypress boards. The insulating properties or this wood are excellent, and the dimensions or the box allowed tor a volume ot water that was so large that the temperature or the water could be controlled vith ease. Stirring in the bath vas provided tor by a turbine type stirrer located in one corner. Tempera ture control was accomplished by an arrangement whereby a portion or the water was circulated by means or a small circulatory pump through copper coils that were immersed in an ice bath. The proportion of' water passing through the cooling coil 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 Beckman pH meter,, model G, vas used f'or all pH measurements in the course or the investigation. In the pH

PAGE 32

... 30 range below s.5 the normal glass electrode"was employed) but at higher 'pH values the measurements were made using a high:, alkalinity, glass electrode. BueriJn!nta} Procgdures apcedYDJ! t2r." preparation .st, werking 12lstion1. Tb.is procedure vas complicated by the necessity tor keeping oxygen out or oontact with the solutions and the various reagents as the •olutionswere 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 tirst steps in the procedure were to f":i.11 the •eries or reaction vessels with diatilledn.ter 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 or gum rubber tubing and projecting through 1-hole rubber stoppers were then inserted into the larger or the tvo holes 1n each or the lids so that they extended to the bottoms or the jars. The smaller of" t.he two holes in each or 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 or nitrogen through the mercury in the pressure regulator. In the ease of each vessel in turn the stopper was remoYed trom the smaller hole in the lid., the water was torced out and replaced by nitrogen., and the stopper

PAGE 33

31 . was r~placed. Mter all the jars had been treated in this manner ~ stoppers were again removed; and a stream of nitrogen vas allowed to . tlov through all of" them simultane ously . until . no more than a : drop or . two of' water remained in ea _ ch . vessel., The -v:essels were then restoppered, and the flow of nitrogen through the pressure regulator . was reduced toa slow bubbling . The next steps in . the < procedure deal : vi th the prepara t:1.on <. of' . the working butf'er solu~ions in .the . nitrogen ~illed reaction vessels; and . they were completed for each vessel once they ' vere begun . , A jar vas removed . f'rom the rack, the nitrogen tube was removed, and o. stopper vas immediately put in its place. Now the veasel ttl\S , placed in a position under the : diluting uater pipet (see B, Figure 1), : the stopper vas removed trom the smaller of' the two . holes in the lid# and a nitrogen tube located near the pipet was immediately inserted to a , : depth or about one inch into the jar. After removal of' : the other stopper f'rom the lid the leg or the pipet was inserted, and the jar was raised until the tip or the pipet leg vas at tho bottom of'the : veasel. 500 ml. of' the oxj'gen.;.f"ree diluting vater were measured into ~he vessel f'rom ~he pipet vhile a stream or nitrogen from the nitrogen tube prevented air .from entering. Next the jar vas lowered from the pipet and placed on ~he desk top, the stream of nitrogen still flowing f"rom the nitrogen tube. . 25 . ml . of' ox7gen-tree, concentrated, but"ter aolution vere measured in b7 means of a ptpet, and atter r~

PAGE 34

32 moval or the nitrogen tube the vessel was restoppered. The jar vas swirled several times to insure thorough mixing. and then it vas placed in the thermostat,. where it was lett f'or two hour• to attain the desired temperature for the subaequent reaction.; After all or the reaction vessels of' the series had been subjected to the treatment outlined in tho above paragraphs .. they were ready ror the introduction ot the stock suJ.:ride solutions andehlorine water. The addition ot ehlorine'water :rollowed very. oloseiythe addition of' the stock sulf'idesolu tion in eaehYessel. However,. in the 1.nterest,ot saving time a schedule. governing these. steps in the procedure was drawn up tor each experiment., and in 'accordance with this schedule the reactions in some of' the vessels ot a series were com pleted and the results determined before the additions were even made to other vessels or the series. Consequently. the procedure for the additions is outlined aa it was followed tor an individual jar. The jar was removed f'romtbe thermo stat., the smaller hole in the lid tmstoppered and the nitrogen tube inserted. Vith a gentle stream or nitrogen tlonng through the tube., the other stopper vas removed from the lid, and the tip of the stock sulride solution buret (see D., Figure 2) was.tnserted to a point beneath the surface or the butter solution in th0 V6$Sel •. The volume or stock sulfide solution required to give the desired sulfide concentration., as ca.lcu lated :rrom the daily standardization of the solution., was

PAGE 35

33 measured into the vessel from.the buret. The buret tip vas 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 Hor I, Figure 2); the size of the buret depending upon.the calculated volume to be added, was inserted beneath the surface of' the solution; and the vo-1ume to be added., as calculated from the daily standardization 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 stdrling rapidly in smnll circles 12 15 times. The jar 1-ras replaced in the_ thermostat to await the previously decided upon time tor stopping the re action and determining its extent. The total time elapsed. while the jar was out ot the thermostat tor the stock.sulfide solution and chlorine water additions was 2 3 minutes. Analxtiea} methgds. After the reaction time that had been previously decided upon for n particular vessel had elapsed,. the extent of' the reaction 1n that vessel"was de termined. 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 otthe reaction. However, in the borderline

PAGE 36

34 cases where it was ditNcult to predict what the residual would be the assumption was made that 1 t vas chlorine, and the procedure f'or the determination or residual chlorine was followed to the point where iodine either -was or was not liberated t"rom potassium iodide in the acid solution. Close observation at this point in the procedure shoved whether the analysis was to ~e made tor residual chlorine or re sidual sulf'ide. The analytical method used tor the quantitative I determination or residual chlorine in the investigation vas the standard iodometric method tor the determination of' chlorine invater (50). The reaction in the vessel was stopped at the end of' the reaction time by the addition of' o.75 grams or potassium iodide dissolved in enough acetic acid solution to lover the pH or the working solution to a value between 3 and 4. Usually 10 ml. of the acid was su:r f'icient. After a thorough mixing by swirling., the solution was poured rrom the reaction vessel into an 800 ml. beaker., and the vesael was washed with a few portions or distilled water from a wash bottle. The liberated iodine was titrated with the standardized 0.01 N sodium thiosultate solution., using 5 ml. or starch solution as the indicator. The quantity or residual chlorine was calculated rrom the titration. The method used tor the quantitative determination or residual sulfide was baaed on the standard io4imetr1c method

PAGE 37

35 for.the determination ot sulfides in water (51). Enough acetic aoi~ solution was added to lower the pH value ot the solution to 3 or 4, and immediately thereafter an excess.or 0.01 N iodine was measured into the reaction vessel by meana or a pipet. These additions were made to the reaction vessel through one of the holes drilled in the lid, and.every pre caution vas taken to prevent the loss ot hydrogen sulfide trom the Jar. ,. Af'ter thorough mixing the solution was poured into a.n 800 ml. beaker, and the reaction vessel was washed with several portions of distilled water trom a wash bottle. The excess iodine was titrated with the standardised 0.01 N sodium thiosultate solution, using 5 ml. 01' 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 .th1osultate solution. The amount ot' residual sulfide was determined trom the amount of iodine that reacts•" Bxuriments lSt. deterrnl.Pe ,a:ept at e9ne1n\rati;ons. The concentrations of sulfide selected tor these experiments were tliose obtained by the addition or 1~ a and 3 milligrams ot sult'ide to the working butter solutions, these additions yielding concentrations or 1.9_ 3.,.8 and 5.6 parts per million as sulfide or a.o# 4.0 and 6.0 part• per million as hydrogen sulfide, respectively. Separate experiments were conducted at each sulfide concentration, and each experiment was made

PAGE 38

36 with a series of seven or eight reaction vessels. Pive ot the vessels received. the care~11 measured ~mount of' sulfide in accordance Yith,the procedur, previously described f'or tlu preparation or the working solutions. and the remaining vessels of' the seri~u, were designated as blanks., receiving no sulfide. The chlorine dosages added to the five vessels con t~ining the sulfide solutions in each experiment were 2, 4, 6, 8 and 10 times the sulf'ide dosage tor ~e experiment. In the case or the vessels designated as blanks, chlorine dosages •ere added that were designed to give chlorine.re siduals or the same order of' magnitude as those in the other jars or the series. Experiments dealing with each of' the selected sultide,ooncentrations were made ai pH values 5.0, 6.o, 7.o, s.o and 9.0. 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 f'rom the previously described procedure f'or the pr~paration of' working solutions occurred in the ease of' the pB so lutions. In this particular case 25 ml. or the diluting water were withdrawn with a volumetric pipet f'rom each of' the reaction vessels 1 and 50 ml. of' the concentrated btf'er . solution were added to the remaining 475 mi thus making a total volume of' 525 ml. of' working butf'er,solutionas :ln the other cases. beerimentf ,!a detrmine !Dit nrecisten st A dtt1mtrn1t;L9n. In these experiments each reaction vessel in a series of' f'our

PAGE 39

37 .Tessels was treated in a manner as nearly identical aspossible with the treatment received by all the other vessels of the series •.. Two milligrams of sulfide were accurately measured into each or the vessels or a series in accordance with the procedure for the.preparation or the working solutions;. A chlorine.dosage or 17.68milligrams was added in each case, and the-reaction time and the temperature were held conatant at 20mtnutesand 25 c., respectively., for all the experi ments. The experiments were rwint pH values 5.o, 6.o, 7.0 and: a.o. The deviation previously mentioned in eonneetion. with the•preparation of the working buffer solution at pH 5 also occurred duringthese experiments •. , ' IIRfltWPt! j,a, dtfrt?r.mint .Ya. e1t1nt at immtfiai1 tt actiS!P• These experiments were designed to give an indication or just how tar the reaction progresses at'the instant the re• aetants are brought into eonta.ct with each other. 0.75.gram or potassium iodide was dissolved in 5 ml. of' starcl1 indicator solution in each one of' tour 800 ml. beakers. Working butrer aolutionswere prepared in each of' a series of fourreaetion vessels 1n.aeeordance with.the procedure previously outlined. The8e butter solutions, each in ~ts turn# were earef'ully poured into the 800 ml. beakers in such a manner that the dis solution of oxygenfrom the air vas at a minimum, and.four milligrams of' sulfide were accurately measured beneath the surface.of the solution. The tip of' the 10.;.m1. chlorine water buret wasin•erted beneath the surf'ace of' the solution, and

PAGE 40

38 the sul:fide solution was immediately titrated with the chlorine water until a blue tinge appe.ared in the aolution and persist ed tor a tev seoonds. This was taken as the end-point or the titration., the potassium iodide acting as an oxidation-re duction indicator in the presence of' the starch. The experi ment vas conducted at pH values 5.0., 6.0., 7.0., s.o and 9.0, and the temperature was held constant at 25 c. tor all runs. The deviation that has been mentioned concerning the prepa ration or the pB 5 working hurter solution occurred .again dur ing these experiments. 1xum:1mn .12. determine the 1t'lgct at .lla• The concen tration of' sulfide employed in these experiments was 3.8 parts per million as sul:fide or 4.0 parts per million expressed as hydrogen sulfide., which was the concentration that resulted when two milligrams or sulfide were added to the working butter solution. Each experiment was made Vi th a aeries or eight reaction vesaels 1 two or which served as blanks contain ing no sulfide. Each or the remaining six Teasels received two milligrams or sulfide in accordance with the procedure tor preparing the working solutions., and then a dosage or 17.68 milligrams or chlorine vas measured into each one in its turn. The chlorine dosage added to the blanks was one that was calculated to 7ield a chlorine residual comparable to those in the other six reaction vessels. The residual chlorine in each ot the .six vessels that were dosed with sul :fide ns determined at the end of 1,, 5, 10., 20, 40 and 80

PAGE 41

39 minutes., respectively. The same reaction times were employed 1n . the _ case or the blank solutions. However, in order to_ get _ a complete, blank run .of' six determinations covering . the .. entire range ot ~eaotion , times employed in the_ experiments, _it va_s necessary to .oombine blank determinations _ror .t~eruns._ at the_aame pH value. Experiments were run at,Pf:l Y~lues 5.;J, 6.0, 7.0,. s.o and 9.0* while the temperature in all of-the ' . . . experiments.was heldconstantat25 c. The deviation occurring in ._ the prepare. tion of", the pH . 5 working buf'f"er . so lution .is again pointed out in connection with these experi ments •. Ruer:&m!ntp l2. .c1ei12trnim! .l!!I. e[ttet .at te,mamtm:•• Ex periments. identical to those described in-the above paragraph were run atpHv~lues 5.0,6.0# 7.0, s.o and 9.o, but temper atures of' 15-and 20 c. were substituted at each pH Yalue tor the 25 9 c. ,mplpyed 1n the.previous runs. Bu,riment,s a determtne M!I. eueet at i211ic 1tr,yth. These experime11ts were run 1n a manner similar to that used to .determine th& ef'.teet of" time on the reaction •. However, provision vas made in the procedure.for changing the ionic strength value of.the working solutions f'rom one.experiment to:anoth~r# and t'urthermore all the experiments were run at a constant pll value or 7.0" as nearly as the changing ionic strength would pe~t. A previously deaer~bed potassium chloride solution was used to vary the ionic strength. The substitution or this potassium chloride solution ror dilut

PAGE 42

40 1ng va:ter in the reaction vessels caused a deviation in the normal procedure tor the preparation ot the working solutions. The potassium chloride solution was substituted by 10 ml. inorementsJ therefore_ the diluting water was removed b7 10 ml. increments with volumetric pipets in order to make room in the solution ~or the chloride solution so that the final vo.lume ot 5~ ml. tor the working butrer solution would re main unchanged. The ionic strength vas .varied trom 0.05463 to o.2046,during.the experiments-by substituting volumea or potassium chlo~ide aoluticm f'rom o to 60 ml. in this ~ashion •. E!iP!!rimftnts !a det1tm1Jtt _ the ettgot s.t ehlpfidtl• The chloride expe,riments :we~e run in the same way that the ionic strength experiments were run. with the exception that the devi.ations trom. the. normal. procedure tor. the preparation ot the working solutions were employed to var1-the chloride concentrations from one experiment to another. leaving the ionic strength value constant at 0.1093. Allot the experi ments were run at a pH value or 7.0 and a temperature or 25o C, In order to vary the chloride oo_noentration and .leave the .ionic strength value constant in the v:orking butters it was necessar1 to substitute potassium chloride solution of' a stated strength tor concentrated butter solution in the pre paration.of the working butter solutions. The_ potassium chloride solution used tor this purpose has been described. In that experiment where no chloride ion was added., 50 ml. portions ot concentrated butter solution (pH 7.0) were added

PAGE 43

41 to 475 ml. portions or diluting water obtained by withdraw-. ing 25 ml. from each reaction vessel by means or a volumetric pipet. This resulted 1n a working buffer solution having an ionic strength value of 0.1093. In the other experiments thi8 same ionic strength value was maintained., but the chloride concentration was increased by regular 1ncreme11ts by .the substitution of' 5 ml. increments of' the potaseium chloride solution tor equal portions of' the concentrated butter solution 1n the working solutions. During the course o.r these experiments the chloride concentration was varied .from .o to about 2,000 parts per million by increasing the volume o.f the potassium ehloride in the working butf"er trom. Oto 25 ml. and decreasing the concentrated buf'f'er from 50 to 25 ml., the respe~tive increases and decreases being made by 5 ml. increments~ ID!tWUll lS!. dgt!U!1!1iJ!e the. ett19t s.t J;a:d;cpQD-1,on egncentrat&99. No additional procedures were required in the study of' the etteet of hydrogen-ion concentration on the oxidation. The procedures used in theexperiments con cerned with concentrations and time were designed to lend i~f'ormation at the same time to the study of the ettect of' the byd~ogen-ion concentration.

PAGE 44

EXPERIMENTAL DATA The results of the various experiments conducted during this investigation are expressed as ratios that show the number of units of chlorine reacted per unit of sulfide re acted. These ratios are given both 1n units of' milligrams and moles. The milligram units are given because it is easier for one to realize in milligrams than in moles the quantity of chlorine required to react with the sulf'ide, and it is a relat_ively simple process to calculate dosages tor comparative purposes from data of' this type. The ratios are presented in mole units tor two reasons. In the f'irst 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 tree sulf'ur 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: Cl 2 /. s•2c1• /. S 4Cl 2 /. s-/. 4B 2 0: ~HCl /. so 4 -In accordance With these relationships the respective ratios for the formation or tree sulf'ur and sulf'ate are the rollowing: Mole1 of' chlorine : 1 Mole of' sul.f'ide Moles of chlorine = 4 Mole of' sulride 42

PAGE 45

43 Since the products are tormed simultaneously the ratio Will generally assume some value between one and tour. Therefore. considering that the ultimate end-product ot the oxidation under the most favorable conditions must be sulfate alone. the magni.tude or the value or the ratio may be taken as an indication ot the progress or the reaction. The larger the ratio the greater is the quantity or sulfate that results from the oxidation. The other reason tor the use ot the mole as the unit in expressing the ratio deals with a simple method tor determin ing the amounts ot tree sulfur and sulfate formed by the re action. Choppin and Faulkenberry (46) have indicated a method whereby the amounts or sulfur and sulfate produced can be calculated trom a knowledge or the ratio 1n this form and the total quantity ot sulfide oxidised. The method in volves the use or two simultaneous equations developed from the above-mentioned stoichiometric relationships concerning chlorine and sulfide. Starting with one mole or sulfide and letting X mole go to sulfur and Y mole go to sulta~e. one can obtain the equations, X /. Y = 1 (with respect to sulfide) and X /. 4Y: experimental ratio (with respect to chlorine). which can be solved tor X and Yonce the experimental ratio bas been established. The amount ot tree sulfur formed can be calculated in any desired units by multiplying X by the quantity of sulfide oxidised expressed in the same units.

PAGE 46

44 Similarly, the amount ot sulfate can be calculated by multiply ing 2.99Y by the quantity ot sulfide oxid.iaed expressed in the desired units. Table 2 gives_ the values or X and Y calculated from certain values ot the experimental ratio. Also included in the table are values tor the ratio or moles-ot sultur per mole or sulfate produced, which values illustrate the effect iveness or the experimental ratio as a means to indicate the progress or the reaction. The renlts or the ex,eriments designed to indicate the precision to be expected in the overall treatment ot each individual reaction vessel are presented in Table 3. It is seen from these results that the precision or the determi nation of' the experimental ratios in the manner adopted tor this investigation is dependent upon the hydrogen-ion con centration or the solution .in which the oxidation takes place. Much greater precision is indicated in acid solutions than in solutions having pH values larger than 7.0. However, con sidering the numerous time consuming steps and measurements that must be taken in the preparation ot the working solutions, the seant quantities o~ the reactants and the ever preaen~ possibility or some contamination by oxygen from the atmos~ phere, the precision or the determination appears to be satisf'actory throughout the range or hydrogen-ion concen tration included in the investigation.

PAGE 47

TABLE 2 RELATIONSHIPS EXISTING BETWEEN CERTAIN VALtJES FOR THE EXPERI MENTAL RATIO AND THE MOLES OF SULFUR AND SULFATE PRODUCED .PER HOLB OF SULFIDE OXIDIZED. &xperimental Moles Sultur Moles Sulfate HoJ.e! Su:ltur Ratio:; per Mole of' per Hole of' Mole Sulfate Suttide {X) Sylf';lde 'Y) (x{!) 1.00 1.00 0.000 ___ _, 1.10 o.967 0.033 29.3 1.20 0.934 . o.oaa 14.2 1.ao 0.900 0.100 9.00 1.40 . o.aa1 0.133 6.52 1.50 o.saa o.167 5.00 1.60 o.soo 0.200 4~00 1.70 0.767 0.233 3i 1.so o.733 o.267 2.75 1.90 0.100 0.300 2.33 2.00 o.667 o.aaa 2.00 2.10 o.633 0.367 1~73 2.20 0.600 0.400 1.50 2.30 o.567 o.433 1.31 2.40 o.533 o.467 1.14 2.50 0.500 o.5oo 1.00 2.60 o.467 o.533 0.876 2.70 0.433 o.567 0.763 2.80 0.400 0.600 o.as7 2.90 0.367 0.633 o.579 3.00 o.aaa o.667 0~500 3.10 0.300 0.700 0~428 3.20 0.267 o.733 o.364 a.so 0.233 0.767 0.304 3.40 0.200 o.soo o.250 3.50 o.1s1 0.833 0.200 3.60 0.133 0.867 o.153 3.70 0.100 o.Doo 0.111 3.80 0.067 0.933 0.011 3.90 0.033 0.967 0.034 4.00 0.000 1.00 0.000 -.45

PAGE 48

TABLE 3 ~CISION OF DETERMINATIONS AT VARIOUS pH VALUES Volume 'of buffer solution: Sttlf"ide added: Concentration or sulfide solution: Chlorine added t . Reaetion•time: Tempera:ture: '' pH 4.93 ,' ' fl' 7.05 7.92 Chlorine Reacted (mg) .. 15.60 15.55 15.58 12.20 12.09 12.17 10.44 10.10 10.22 s:35 8.51 8.35 8.24 Chlorine to Sulficle Ratio ,(g) 7~80 ,. 7.78 7.79 6.05 5.11 4.18 4.26 4.18 4.12 -~46 2 mg/529 ml ' Chlorine to Sulfide Ratio. (mole) . 3.53; 3.52 .3.52 2.31 1.89 1.93 1.89 1.86 525 ml. 2.00 mg. or 3.78 ppm. 17.68 mg. 20 min. 250 C. " Devi.a t1on . from A.v
PAGE 49

47 ga,ectgt ((gpcentrata.ons.sm tbft Jteac;&un The experiments concerning the e.t.teot o.t the concentra ... tions or the reactants on the oxidation were designed to in dicate the extent of the reaction under various conditions or concentration with the other conditions being held constant. It has been mentioned previously in the description or the procedure .tor these experiments that experiments were run at pH values varying by unit increment• from 5.0 to 9.0 and that the reaction time and the temperature vere held constant tor all experiments. Attention is invited to Table ' . 1 tor information relative to the ionic strength values and . . the chloride concentrations or the working butter solutions in which the oxidations were carried out. . . The results or the experiments are presented in Tables 4 6# inolusi"f'e, which deal with 1.90,. 3.78 and 5.65 part per million aul.tide solutions, respectively. It is seen . from these tables that, within the limits or the experimental . . error involve4 in the procedure, the chlorine to sulfide ratio is independent ot the original concentration or sulfide in the range ot sul.t1de concentrations investigated in the experi ments. Furthermore, it appears that the results obtained vith the 3.78 and 5.65 part per million sulfide solutions are more consistent than those Vith the l. 0 90 part per million solution. This is probably due in part to the greater accuracy that is possible in measuring out the larger volumes ot reagents re

PAGE 50

TABLE 4 BPPECT OF C0NOBKTRATI0NS ON TBE RATIO OF CHLORINE TO SULFIDE REACmD. SULFIDE CONCENTRATION IS 1.90 PPM. Volume of' butter solution: 525 ml. Sultide addedf. 1.00 mg. Concentration ot sulfide solution1 1 mg/527 ml or 1.90 ppm. Reaction 'time: 20 min Temperature: . . 25 c. pH Chlol"ine Chlorine Sulf'ide Chlorine. to Chlorine•to Added Reacted Reacted Sulf'ide Sul.tide (mg) (ms;) (mg) Rat,;!s, 4.88 2.00 2.00 0.318 6.28. 2.84 4.00 4.00 o.574 6.97 3.16 6.00 6.00 0.826. 7~26 3.29 . s.oo 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 ft 4.00 4.00 0.848 4.72 2.14 a.oo .5.06 1.00 5.06. . 2.30 II a.oo 5.56 1.00 5.56 2.52 10.00 6.00 1.00 6.00 . 2.72 7.05 •, 2.00 2.00 0.510 3.92 1.78 4.,00 4.00 o.952 4.20 1.90 6.00 4.82 1.00 4.82 2.1s s.oo 4.76 1.00 4.76 2.16 10.00. 4.93 1.00 4.93 2.23 7.92 2.00 2.00 o.592 3.38 1~53 4~00 --1.00 ----6.00. 4.15 1.00 4.15 1.88 s.oo 4.12 1.00 4.12 1.87 10.00 4.21 1.00 4.21 1.91 8.95 2.00 2.00 o.1>s2. 3.56 1.61 ff 4 •. oo -1.00 ....... ---6.00 -91!' 1.00 ........ " s.oo 4.23 1.00 4.23 l.92 10.00 4.31 1.00 4.31 1.95 48

PAGE 51

TABLE 5 EF.PBCT OF CONCBNTRA.TIONS ON THE RATIO op CHLORINB . TO SULFIDE HEACTBD •. . . . . SULFIDE CONCENTRATION IS 3.78 PPM. Volume of'butter-solution: :525 mi. Sultide addedt . 2.00 mg. Concentration of' sultide solution: 2 mg/529 ml or 3.78 ppm •. Reaction time: 20 min. Temperaturet .. 25o C.: .. pH Chlorine . •Chlorine. Sulfide. Chlorine to Chlorine. -to Added Reacted Reacted Sulf'ide Sulf'ide II. bud (v> .atj,o 4.90 4.00 4.00 o.694 . 5 •. 77 2 ... 62 s;.oo s.oo 1.18 6.78 3.07 12.00 ,12.0(l 7 •. 06 3,.20. 16~00 l4.9!J 2.00 7 •.. 49 3.39 20.00 l6.04L . ' . . 2.00 s •. 02 3.63 5.,89 4.oo 4.0(~ 0.901 4.44 2.01 .. s.oo s.ou 1.61 4 •. 97 2.25 " 12.00 11.04k 2.00 5.52 2.50 16.00 12.lU 2 •. 00 6.05 2.74 20.00 12.8(1 .2.00 6.40 2.90 7.05 4.oo 4.0(1 1.02 3.92 1.78 s.oo s.oc, 1.89 4.24 1.92 12.00 9.5!:il 2.00 4.80 2.1s 16.00 9.9(!. 2.00 4.98 2.26 .20.00 ,. 10.3.fi 2.00 5.17 2.:H 7.92 4.00 4.,0fll 1.18 a.aD. 1.54 ff s.oo .......... 2.00 -... .......... 12.00 8.14: 2.00: 4.07 1.84. 16~00 8.16, 2.00 4.08 1.85 It 20.00 s.ae: 2.00 4.14 1;87 8.95 4.00 4.0(]1 1.13 3.54 1.60 s.oo -2.00 -........ _ 1t 12.00. 7 .• 5fJt 2.00 a.75 1.70 ti 16.00 8.23: 2.00 4.12 1.87 .20.00 . s.a11 2.00 4.14 1.-87 49

PAGE 52

TABLE 6 EP.FEOT OF CONCENTRATIONS ON THE RATIO Oil' CHLORINE . TO SULFIDE REACTED. SULFIDE CONCENTRATION IS 5a65 PPM. Volume of buffer solution.: 525m1 .• Sulfide' addedt ' . a.oo mg. Concentrationor sulride solution: 3 mg/531 ml or 5,. 65 ppm. Reaction time.t 20 min. Temperature: 25o c. pH .Chlorine Chlor:lne Sulfide Chlorine to Chlorine to Added ' ' Reacttad Reacted Sul.tide Sulfide {st (mg} (ul natto (m,e;) !Jattg Cm9J,e} 4.90 6.oo. 6.00 1 .• 01 5.94 2.69 12~00 12~0t) 1,,74 6 .• 90 3.12 " 18~00 1s.oo 2.51 7.17 3,.25 24~00 22.s;a 3.oo 7.61 3.4.5 ao.oo 24.1:? 3.00 8.04 3.64 5.89 6.00 G.oo 1.27 4.72 2.14 " 12;00 12.00 2.29 5.24 2.;37 " 10.00 16.9)L a.oo 5.64 2~55 " 24.00 17.7:t 3.00 5.91 2.68 ao.oo 18~30 a.oo 6.10 2.76 7.05 .6.00 6.on 1.50 4.00 1.a1 12.00 12.00 2.76 4.34 1.97 ft 18.00 14.lfi 3.oo 4.72 2.14 It 24.00 14.Sfi• a.oo 4.95 2~24 It ao.oo 15.61. 3.00 5.20 .36 7.92 s.oo e.oo, 1.74 3.45 1.56 " 12.00 .--a.oo ---ff 1s.oo 12.oe a.oo 4.02 1.82 tf 24.00 12.26 3.00 4.09 1.85 ao.oo 12.65 a.oo 4.22 l.91 6.00 6,.00 1.66 3.62 1.64 ., 12.00 ---a.oo ------ff 18 •. 00 11.94 a.oo 3.98 1~80 ft 24.00 12.35 a.oo 4.12 1-.87 30.00 12.66 3.00 4.22 1.01

PAGE 53

51 quired tor the more concentrated solutions. At any rate the results tor the 3.78 and 5.85 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 or curves 1n this figure. Probably the most striking feature illustrated is the clear cut dependence or the experimental ratio upon the hydrogen ion concentration or the reaction medium. 'ftle ratios were quite large 1n a reaction meclium having a pH value or 4.90, :t.ndicating that the larger portion or the oxidised sulfide was o:xidiZed all the way to•sulfate. However, the ratios de creased 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 or s.o. The figure also brings into sharp focus the ettect ot the ratio or chlorine added to the sulfide on the experimental ratio and the dependence or the magnitude or this ettect on the hydrogen-ion concentration or the reaction medium. The ratio or chlorine to sulNde reacted is seen to have increased in all cases nth an inereaae in the ratio or chlorine added. Thia increase vas quite substantial in,aoid solutions~ but it became less marked as the pH value or the medium was increased •. Finally, at a pH value in the neighborhood or s.o the oharac teriatics or the increase in the experimental ratio With in creased chlorine ratios appear to have become fixed and much less impressive than 1n acid solutions.

PAGE 54

TABLE.7 EfflCT OF CONCENTRATIONS ON THE RATIO OF CHLORINE TO SULFIDE REACTED. SUHMART OP EJPBR.IMBNTS . . pH.• ,. '. ' Mole Ratio or , Chlorine Average "Deyiation to Sulfide Reacted Ratio :f"rom Sultide Cone. Sul.tide Cone. Average , a.78 ppm. .65 ppm .. . {l!ble l (l!l!:&1 sU 4.90 2.62 2.69: 2~66 1.3 3.07 3.12 3~10' o.s 3.20 3.25 3~23 o.a 3.39 3.45 3.42 0.9 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 It 2.50 2.55 2.53 1.0 2.74 2.88 2.71 1.1 2.90 2.76 2.83 .5 7.05 1.78 1.81 1.so 0.9 1.92 1.97 1.,95 1.3 2.18 2.14 2~16 1.0 2.26 2.24 2.25 o.5 1t 2.34' 2.36 2.35 o.5 7.92 1.54 1.56 1.55 0.7 ...... _ --.... .. ..... _ -1.84 1.82 1.83 0.6 1.85 1.85 1.85 o.o It 1.87 1.91 1.89 1.1 8~95 l.60 l.64 1;.62 1.3 •' -..... ----1.70 1.so 1.75 .s 1.87 1.87 1.87 o.o .~7 1.91 1.89 1.1 52

PAGE 55

"t:, \J ti Qj ")5 .:;:: 'V) -I: 'lJ t:: ., I.,. --1:: \.J " <;) .I::> ' ....... ti Q( \) '-.. 53 Fl GUR .3 EFFECT OF CONCENTRATIONS ON THE.. RATIO OF CHLORINE TO SULFIDE REACTED. AVERAGE RATIOS FOP. 3.78 ~J 5.G5 ?riv!. Sl/LFIDE SOL!JTIONS PLOTTED. 3.7. 3.SO 3.30 3.10 2.'?JO 2.70 2.50 2.30 2.10 /.90 1.70 1.50 O ( See Tobie 7) 0 pl/= 7.~2(0) p/1.95(,!j) 2 4 , 8 Rolio of Chlort"ne Added fo Svlf'ide M_g.

PAGE 56

54 Figure 3 capably illustrates another important consider ation. It is noted trom Tables 4 6., 1nclus1•e., that in every case the quantities of chlorine that must be added to completely eljmjnate sulfide from the solutions are greater than the stoiolrl.ometric quantity-require4 tor oxidation to tree sultur., and these quantities are observed to depend to a great extent upon the hydrogen-ion concentraUon of the solution. Thus it :ls indicated that there is an oxidation process whereby the oxidation of a portion of the sulfide is carried beyond the tree sultur atage while there is yet unoxidized sulfide present 1n the solution. This is graphically shown in Figure 3., where 1.t may be obsened that the exper1mental ratio is always considerably greater than un:lt7 even when the chlorine is ad4ed on a mole to mole basis (2.2 milligrams per milligram or sulfide) or slightly less. Prom Tables 4. 5 and 6 it is determined that the chlorine dosages required to insure the complete elimination or unoxidised sulfide 1n reaction media of various pH values lie between the following •aluesr pH Chlorine Dosage 4 .. 90 5,.89 7.05 7.92 S.95 W Bf mg. ot <&de) 6 8 4-6 4-6 2-4 2-4

PAGE 57

55 It became eTident from preliminary experiments concern ing the rate of oxidation of sulfide solutions b7 chlorine that the initial oxidation process is a very rapid one. This initial portion of the oxidation was found to giTe var to a much slower process. The experiments shoved that the more rapid portion or the oxidation takes place within the first minute of the reaetion 1 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 encl of this reaction time was always considerably greater than unity• indicating that the oxidation at that point had proceeded beyond the sultur state. Inasmuch as previous investigators have suggested that sulfur is the primary product ot the oxidation . . (46)., it was decided that the study or the oxidation rate could best be made in two parts. The first part or the study was to deal with the initial or rapid oxidation process that takes place tdthin the first minute or the reaction. The procedure tor 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 chlorine water under various known conditions, using potassium iodide in the presence of a little starch as an oxidation-reduction type indicator.

PAGE 58

56 This particular indicator was chosen because a study or the pertinent oxidation potentials reveals that it s~ould indicate in acid solutions when the oxidation or the aultide to sulfur has just beeri completed, and 'thus it should be possible. to determine by its use whether or ~ot it is pl'"obable that the immedi~te'reaction involves oxidation or the'sultide to f'ree ' ' sulf'ur. The literature discloses.that.the above.conclusion in regard t'o the usef"ulne:88 of p~tassium iodide as an :lnd:l.cat or for this particular, titration' in acid solution's has been substantiated by experiment (52). Further co~sid~ration or . . . the problem indicatett'that the lodide~iodine couple.cannot be auocesstully employed m the titration or alkaline sulfide solutions with chlorine to indicate when the oxidation to sulfur is just completed. However, the experiments were con ducted 1n the same manner.in alkaline solutions w1.th the thought that the results would still indicate the extent ot the immediate reaction under such circumstances. The results or the experiments are presented in Tables. It is evident trom these results that the product ot the initial oxidation process in acid solutions is tree suitur. Furthermore, it is just as evident that in alkaline solutions the immediate oxidation process carries the reaction beyond the tree aultur stage, the extent ot the reaction depending upon the alkalinity or the medium within the limits chosen for this investigation. The second part or the study or the effect or time on

PAGE 59

TABLE 8 EXTENT OF THE IMMSDIATE REACTION BBTWBEN CBLORDE .AND.SULFIDE. Volume or burrer solution: 525 ml. Sulfide added: 4.00 mg~ Concentration or sulfide solution: 4 mg/533 ml or 7 .,51 ppm. Temperature: 250 c .. Reaction time:Sulfide solution is titrated vith:chlorine water, using starch and KI as the indicator. pH Chlorine Chlorine to Chlorine to Deviation -Reacted Sulride Sulf"ide . ,from (mg} Ratio (mgl _ Ratj.o (moJ.1) Average_. 4.93 9.34 2.34 1.06 -0.5 It -9.31 2.33 1.06 -0.5 It 9.48 2.37 1.07 0.5 1h47 2.37. la!!Z ().5 Average 1.065 5.89 8.86 2.22 1.01 o.o It s.s7 2.22. 1.01 o.o " s.s9 2.22 1.01 o.o .. 8.89 2.22 1.01 . 'o.o Average 1.01 < 7.05 9.62 2.41 1.09 -o.9 " 9 •. 77 2.44 1.10 o.o 9.83 2.46 1.11 0.9 " 9.84 2.46 lall 0.9 Average 1.10 7.92 11.17 2.79 1.26 o.o 11.07 2 .. 77 1.25 -o.s a 11.13 2.78 1.26 o.o 11 11.23 2.81 la2Z o.s Average 1.26 s.95 12.43 3.11 1.41 1. ft 12.24 3.06 1.39 . o.o 12.10 3.03 l.J! -1.4 Average 1.39 57-

PAGE 60

58 the oxidation was concerned with the reactian that takes place in the reaction mixture af"ter the initial minute of' reaction . ' ' . time. The procedure for the experiments that were inYented to determine.this ef'fect has been described earlier. The sulfide . . . . . conoentratio!1 or 3.8 parts per million was selected f'or these experiments beoaus'e it lies in a range or concentrations f"requently eneounteredin natural uaters, and it involved convenient quantities or reagents. Furthermore, it was demonstrated by the previously mentioned experiments concern ing the erf'ect or concentrations that under constant-conditions the experimental ratio is independentof".the originalsulride concentration as long as the latter is within the concentration 11.J#ts considered in this investigation. The choiceof' 17.68 milligrams was adopted as the chlorine dosage to be added be cause it repre.sents the stoichiometric quantity or chlorine tha:~ would ondiae the two milligrams of'. sulf'ide that were pre'sent 1n. eacb of' the solutions all _the va.y to sulf'ate if' the ~ther conditions of' the experiment would allow it. Further more. the earlier experiments dealing with concentrations show.,. ed that the-chlorine dosage required to insure complete removal of" unoxidized sulf'ide in a reaction medium or pH value_around 5.o: 11es somewhere in the vicinity of" 8 milligrams of', chlorine f .-~ ,4 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 buff'er solutions in• volYed in these experiments.

PAGE 61

59 The results ot the experiments are presented in Table 9. ' It has been to1md trom these results that when t.he common logarithm or the experimental ratio is plotted against the common logarithm or the reaction time a straight line is obtained. The evidence is presented 1n P:l.gure 4, where t.he plots are shown tor all or the experiments. The equations tor these plots are or the type log R: log a J blog t, or R: atb~ •here!! is the experimental ratio in mole units~! is the reaction time in minutes and,,!. and]!. are constants. It is seen that the oons-tant J!. indicates the extent ot the oxidation ' at the end or one minute ot the reaction, and the constant Jt determines the slope ot the line, or in other words it indi.. catea the rate or the change in the experimental ratio with time, Thus the constant l?, may be regarded as an empirical rate constant. The constants derived f'rom the results oteach of the experiments are presented in Table io, and it ia observed that they are dependent to a very large degree upon the conditions under which the oxidation takes place. The ex.per:1• mental ratios calculated by the use ot the constants are also exhibited in Table 10 1 and they are seen to be in good agree ment with the observed ratios. Effect or Temerature on the Reaction The range ot temperatures selected tor use in the expert

PAGE 62

TABLE 9 EPPBCT OF TIME ON THE RATIO OF CHLORINE TO SULFIDE RBAC'fflD. OBSERVED DATA Volume or buf'ter solution: 525 ml. Sulfide added: ' ' ' ' . , 2.00 m,g. Concentration of sulfide solution: 2 mg/529 ml or 3. 78 ppm. Chlorine added: . 17.68 mg. Temperature: 25o a. pH Reaction Chlorine Chlorine to Chlorine to Logarithm Time Reacted Sult:lde Sulfide of Ratio (min} . 'Cul Ra:t,j.o (u) Rat&o (mole) (Mo;&t) 4.90 l 14.92 7.46 3~38 o.529 5 15.05 7.53 3.41 o.saa 1f 10 15.27 7.64 3.46 o.539 .20 15.30 7.65 3.47 ,0.540 " 40 15.46 7.73 3.50 0.544 ff 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 .10 0.431 ti 10 11.96 5.98 2.71 0.433 20 12.20 6.10 2.76 0.441 " 40 12.38 6.19 2.so 0.447 80 12.50 6.25 2.83 0.452. 7.05 1 9.72 4.88 2.20 0.342 ft ,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 2.46 0.,391 It 80 11.22 5.61 2.54 o.405 7.92 l 8.24 4.12 . 1.87 0.272 5 8.39 4.20 1-.90 0.279 10 8.45 4.23 1.92 o.2sa If 20 8.38 4.19 1.90 0.219 " 40 8.75 4.38 1.99 o.299 80 8.70 4.35 1.97 0.295 8.95 1 7.60 a.so 1.72 o.aae tt 5 7.99 4.00 1.81 o.25s 10 8.09 . 4.05 1.83 o.263 20 8.26 4.13 1.87 0.212 It 40 8.41 4.21 1.91 o.2s1 ft 80 8.53 4.27 1.93 0.286 60

PAGE 63

61 FIGURE 4 EFFECT OF TIME ON THE RATIO OF CHLORINE TO SfllF/OE l?EACTEO (See Tobie ) 0.55 0.50 " I:, {; ' 04C -~ ' ,;:, .<::l 0.35 '~ ' 0.30 0 ,;:, t , ... pf!: 7. 92 I:, '-I 0.25 0.2. 0.10.8 1,0 I. 2 I. G I. 8 2..0 Logorilhm of Time -M/n

PAGE 64

TABLE 10 EPP.EC! OP TIME OX THE RATIO OF CHLORINE TO SULFIDE REACTED. pH ., 4~90 ff ff fl ft It " 8~95 tt ff fl CALCtJLATBD DATA . , Source or data for caloulations:bTable 9. Basis for calculations: R: at, where. R = Chlorine to sulfide ratio (mole),. .1: Reaction time (min), : ,!. and hare characteritrt.ic constants. Reaction Constant Constant Time s. ll (min) l 10 20 40 80 1 5 10 ao 40 80 1 5 10 20 40 80 1 5 10 20 40 80 1 5 10 20 40 80 3.39 1.74 tt 11 0.00720 " 0.0193 0.0497 " 0.0124 " tt 0.0243 " 62 Calculated % Deviation Chlorine to f'rom.Observed Sulfide Ratio Rat~o
PAGE 65

63 ments planned to illustrate the effect of temperature on the oxidation extends from 150 to 25 8 centigrade. Th.is range was chosen with the thought that it includes the temperatures o:f the great majority of" sulf"ide-bearing::vaters. ~ther4'.' more,, thttse temperatures were :round to be the moateonvenient . ,. . . ones With which to work, using the equipment at hand. Tem, ' "' perature control with the thermostat constructed :for use in . . this investigation was :found to present a rather difficult problem vhen temperatures below 15 6 or in excess of 25o vere , .. .;• . . . employed. Also to be considered wns the p~obability __ ~t temperatures much greater than 25 degreei(were likely to yield doubtf"ul results because of the advo~se effect o~ the solubilities o:f tbegaseous'reactants. It has been me~t1oneo that tm procedure for the experiments us~d to dete~e the temperature ettect was identical to that for determini~ the efteet of' time on t~e reaction. The time_experiment~ !ere just repeated.at 15 and 20 c., and the ~esults were ~mparoc vith the 25 degree results. The data observed during the 15 and 20 degree experiments are presented in Tables 11 and 12., reapec~ivel:,, ana. t~ey are summarized in Table 13 along with the datf:i :from Table~., which deals with identical experiments conducte~ at 25 degre•s• A s!,udy of' the summary of the., data reveals i:io pattern of: change in the experimental ratios that can be at~ributed t.o t(le variations in temperature. Indeed., in tfu? large majority of' cases the maximum deviations between the ratios for the vari

PAGE 66

TABLE 11 EFFECT OP TBHPERATURE ON THE RATIO OF CHLORDB TO SULFIDE REACTED. OBSERVED DATA Volume of buf1"er solution: 525 ml. Sulfide added: 2.00 mg. Concentration of sulfide solution: 2 mg/529 ml or 3.78 ppm. Chlorine added: 17.68 mg. Temperature: 15 6 c •. pH Reaction Chlorine Chlorine to Chlorine to Time Reacted Sulfide Sulfide (min) (mg) Rat19 Cmg) Batiqn .{moJre) 4.93 l 14.76 7.38 3.34 5 15.17 7.59 3.44 10 15.24 7.62 3.45 20 15.48 7.74 3.50 40 15.56 .7~78 3.52 It so 15.68 7.84 3.55 5.97 l 11.36 5.68 2.57 5 12.05 .6.03 2.73 10 12.02 6.01 2.72 20 12.10 6.05 2.74 40 12.38 6.19 2.so 80 12.43 6.21 2.s2 7.14 1 9.41 4.71 2.14 5 9.93 4.97 2.20 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 s.02 1 8.14 4.07 1.84 5 8.23 4.12 l.87 10 8.52 4.26 1.93 ... 20 8.42 4.21 1.91 40 s.64 4.32 1.96 80 8.90 4.45 2.02 9.15 1 7.06 3.53 1.60 ft 5 7.76 3.88 1.76 n 10 7.95 3.98 1.so 20 7.99 4.00 1.81 " 40 8.17 4.09 1~85 80 8.33 4.17 la89 64

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TABLE 12 BF.FECT OF 'mMPERATURE ON' THE RATIO OP CHLORINE TO SULFIDB REACTED. OBSERVED DATA Volume of' buf"f'er solution: Sulfide addedt Concentration or sulf'ide solution: 2 mg/529 ml Chlorine addedt Temperature: " pH 5.92' .. Reaction Time {pn) 1 5 10 20 40 80 1 5 10 20 40 80 1 5 10 20 40 80 1 5 10 20 40 80 l 5 10 20 40 80 Chlorine Reacted (mg) 14.96 15.05 15.27 15.40 15..;54 15.33 11.35 12.02 11.98 12.12 12.34 12.54 9.83 9.92 10.19 10.44 10.75 10.50 8.29 8.51 s.5s 8.71 8.95 9.15 7.25 7.75 7.92 8.03 s.20 8.29 Chlorine to Sulf'ide Ratio (u) 7.48 7.53 7.64 7.70 7.77 7.67 5.68 6.01 5.99 6.06 6.17 6.27 4.92 4.96 5.10 5.22 5.38 5.75 4.15 4.26 4.29 4.36 4.48 4.58 3.63 3.88 3.96 4.02 4.10 4.15 65 525 ml. 2.00 mg. or a.78 ppm. 17.68 mg. 20 Cti Chlorine to Sulf"ide Ratj.o (mole} 3.39 3.4.1 3.46 3.49 3.52 3.48 2.57 . 2.72 2.71 2.75 2.so 2.84 2.23 2.2f> 2.31 2.36 2.44 2.60 1.ss 1.93 1.94 1.98 2.03 2.os 1.64 1.76 1.79 1.82 1.86 1,ss

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TABLE 13 ~CT OP TSMPBRATORE ON THE RATIO OF CHLORINE 'fO SULFIDE REACT.BB. SlJMMARY OF OBSERVED DATA Source or data: Tables 9• 11 and 12. , pH• Reaction Chlorine to Sul~ide Ratios ,Maximum Time (mole) -~
PAGE 69

67 ous temperatures are probably well within the limits of' experimental error for the methoda employed 1n this investt gati.on. Consequently it is belfeYed ~hat a.ny temperature etteet within the range or temperature used is so small that '. it cannot be determined by th.e 1:n:vestigative procedures . ai,plied,and it is therel"oreconaidered to be inaigniticant from the standpoint :ot thf:s atwty . 1tttct . or HfdCUtn-ion C-oncentmtign gn the lteae;tton . . . . . . . . . . It has been stated in the secUon treating of < expert' . mental methode that no ~ddittonal procedures were required 'in connection with the stud,dealing With the ef'tect of' ~e : 11y~gen-ion concentration on the o•idation_ since those : procedures uaed in the experimenta eoncerning concentrations and time vere . designed to 7ield inf'ormation a( the same t,ime _ _ about the hydrogen-ion ~rtect, ~t the hydrogen-ton coneen 'trat1on has a very decided ettect upon the reaction is ciea.-~ 1:, evident from anexam~ation of' ~he results of" all the _ experiments presented thus tar~ However, in regard to t);iiiJ study on:ly a port:ion of' : the previou data was considered~ . the etf"ect ot the b.7drogen-:len concentration on the . ex: tent or the immediate ..-act:lon has alread7 been noted, and f'or a re-view: or the results or the : experitnents ' on this part ot the . investigation attention ia directed to Table 8 -, The remaining portion of' the study concerns the htdrogen-ion ef'f'eet 1n the reaction mixture after the reaction 11a• been

PAGE 70

68 allowed to progress f'or one minute and longer. in this re .gs.rd particular consideration is invited to the data appear ing in Table 9, the graphical representation of' these data in Figure 4 and the related calculated data that are exhibit• ed in Table 10. The essential facts may be quickly noted by ref'erring to Figure 4. Here it is observed that the inter cepts of' the Log Ratio vs Log Time plots increase to a very pronounced degree with decreased pH value, which means that tbe extent of' the reaction at the end of' one minute varies considerably with the hydrogen-ion concentration. It is f'urther obsened that the slopes of' the plots vary vith the pH value 1n such a manner that there is a maximum slope in dicated in the vicinity of' pH 7.0. Since this slope is in dicative of' tm reaction rate, it is seen that there ts a .Possibility of' correlating the reaction rate with the hydrogen-ion conoentrat~en. As a result of' the above ebservations it was decided that the most satisfactory war to attack this phase of' the study would be to demonstrate the ef'fect of' the hydrogen-ion concentration on the constants ,e. and l?. of' the empirical equation,. R: atb,. which has been f'ound to be applicable to the reaction af'ter one minute or reaction time. The signif'icanoe of' these constants has been pointed out in the discussion of' the effect of' time on the reaction. Experiments such asthose used to determine the ef'f'eet or time on tbe reaction were round to be most suitable f'or

PAGE 71

69 producing the necessary data f'or this part.of the invest!gation. In tact the results f'or the time experiments con. . ' '. ducted at pH values 4.90., 5.89, .7.92 and s.95,. which api,ear in Tables 9.and 10, were also employed in this study. The experimental conditions, other than pH value, in all or these experiments were made as.nearly.identical as possible. In addition to the above experiments similar experiments were conducted at.p}J values 6.35,.6.56, 6.79., G.95., 7.25 and 7.44 in an attempt to determine the pH value at which the indioat.. ; . . . .. '. , . ,. .. , ed maximal r~act~on rate occurs. The results of these additional experiments are ottered in Tables 14 and 15., and all of the data in regard to the eonstants .!.. and !1 that are ' , . ', . . . ' . . . . ' . :;. used.in the sttul.y are summarized in Table 16. Figure 5 presents a "Very vivid., graphical illustration or the effect of' the hydrogen-ion concentration on the re... •. , ' 1 action 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 duri~g the first minute or 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.o. The "Value or. the experimental ratio at this minimum is evidently about 1.7, •Moh means that the ratio of sultnr to sulfate produced is about 3 moles to l according to the figures in Table 2.

PAGE 72

TABLE 14 BPFBCT OF HYDROGEN-ION CONCEN'DtlTION ON THB RA.TIO OF CHLORINE TO SULFIDEREACTED. OBSERVED DATA. Volume or bllf'fer solution: 525 ml. Sultide added: .. 2.00 mg. Concentration of sultide solution: 2 m&/529 ml or 3.78 ppm. Chlorine addedt 17.68 mg. Temperature: 25 c •. Ionic strength valuet 0.1093 pH Reaction. Chloride. Chlorine Chlorine to Chlorine to Time Concentration Reacted Sultide Sul.fide Cmin) r . {pRm) <•> Bitto (mg} Ratio Jmtle) 6.35 1 . 1200 10.66 . 5.33 2.42 5 10.99 5.50 2.49 10 11.20 5.Go 2.54 20 It 11.45 5.73 2.60 40 . ll.,96 5.98 . 2.71 80 12.14 6.07 2.75 6.56 l 832 10.07 5.04 . 2.28 5 " 10.81 5.41 2;.45 10 10.84 5.42 2.46 .. 20 11.57 6;.78 2.62 40 11.96 5~98 2.71 80 " . 12.33 6.17 2.so 6.79 l t35 9.96 4.98 2.26 5 10.71 5.36 2.43 1f 10 10.so 5.40 2.45 20 11.46 5.73 2.60 40 " 11.84 5.92 2.68 80 11 12.18 6.09 2.,78 6.95 l 1540 9.52 4.76 2.16 .5 .tt 10.64 5.32 2.41 10 fl 10.39 5.20 2.36 20 10.98 5~49 2.49 40 11.60 5.80 2.63 so, 1t 11.88 5-_94 2.69 Continued 70

PAGE 73

TABLE 14. Continued pH Reaction Chloride Chlorine Chlorine to Chlorine .to Time Concentration Reacted Sulfide Sulf'ide. .tm&9l .. : (ppm) . (ms) Rati2 (m,r) Ratj_o imo&2l 7.25 :1 466 8.93 4.57 ,2.07 5 " 9.60 4.80 2.17, ,ff 10 9.68 4.84 ,2.19 ,, 20 ft 10.06 s.03 2.28 ' 40 . 10.52 5.26, 2.38 80 If 10.92 5.46 2.48 7.44, 1 213 9.00 4.50 2.04 5 9.33 4.67 2.12 1' 10 It 9.20 4.60 2.os t 20 ; ti 9.96 4.9a 2.26 40 10.18 5.09 2.ao 80 10.39 5.20 2.36, ... 71

PAGE 74

TABLE i5 BFF'ECT OF, HYDROGEN-ION. CONCBNTIU.TION ON TUB RATIO OP' . CHLORINE TO SULFIDE REACTED. CALCULATED DA.TA. , . . . %uree of' data f'or calculations: i9'ble 14. lasis ror ealeulationst . R: at, where !l :_ Chlorine to sulf'ide:ra~io (mole), .1 :: Reaction. time (min), . !l. a~d hare characteris_tie constants. pH. , Reaction , Constant Constant Cal.cu.lated Deviation Time .. .,.b Chlorine .to f'rom Obsened (min) . Sulf'ide Ratio Rat;&o ,m21sd 6.35:. 1 2.38 .0.0333 a.as . -1.7' 5 2.51 1.2. 10 2.57. 1.2 " 20 2.63. 1.1: 40 2.69. -1.1 80 2.75 o.o 6.56 1 2.25 0~0502 2.25 -1.a tt 5 2.44 -0.4 10 tt 2.52 2.4' 20 :2.61 -0.4: 40 ft 2.70 -0.4 80 2.80 o.o, e.79: 1 2.24 0.0473 2.24 -o.9 5 tt -2.42 -0.4: 10 2.50 1.0 20 " 2.59 -0.4: 40 2.67 -0~4 80 •,." a 2.76 o.o 6.95 1 2.16 0.0002 2.16 o.o '5 2.34 -2.9 10 " .. 1f 2.43 2.9 20 2.51 o.s 1t 40 .. 2.60 -1.1 80 2.69 o.o Continued 72

PAGE 75

TABLE 15 Contin~ed pH Reaction Constant Constant Calculated "Deviation time a l?. Chlorine to f'rom Observed (min) Sulfide , Ratio Rati.9 (ele) 7.25 1 1.99 o.0499 1.99 -3.9 5 2.16 -o.5 " 10 2.24 2.3 20 2.31 1.3 40 2.40 o.s 80 2.48 o.o 7.44 1 2.03 0.0349 2.03 -o.5 5 1t 2.14 . 0.9 10 " 2.19 5.1 20 2.25 -o.4 40 2.30 o.o 80 2.36 o.o

PAGE 76

TABLE 16 BPPECT OF BYJ>ROGEN-ION COHCBNTRA!lON ON THE CONST.ANTS OF THE EMPIRICAL EQUATION~ R : at SlJMK.ARY OP DAT.A. Source ot data ' : Concentration of.' sulf'ide solutions: Chlorine added: Reaction time: Temperature: pH Ionic Strength Chloride Tables 1, 10, 14 and 15. 2 mg/529 ml or 3.78 ppm. 17.68 mg. 1 so m1ngtes. 25 c. Constant Constant Value Concentration a 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 o .. . 0473 " 6.95 . 0.1093 1540 . 2.16 .. 0.0502 7.2f?, 0.1093 466 ' l.99 o.0499 ,, 7 ... 0.1093 213 , ; 2.03 .. 0.0349 7.9~ 0.2154 0 '; l.87 0.0124 . , s.95 0.1093 3470 1.74 0.0243 74

PAGE 77

75 FIGL/RE. 5 EFFECT OF HYDR06EN-ION CONCENTRATION ON TH CONSTANTS OF THE EMPERICAL Et-;VATION, R = afh. ( See Table /t;,) Curve I: Consfonl a. 3.8 Corve II. : Consfonl-k.. 3.G 0 3.1 3.2 3.0 ~I ..,._ c..B 1/) t: ~I <...> 2.G "> 2'.4 2.2 2.0 /.B L 0 /. C, 1.1, 8 "t. Sc 5. r; G.O ,.4 w.B Z2 7. G, 8.0 8.4 8.8 pH 0.054 o.oso 0.01-~ 0.1)12 IJ.038 0.034 0.030 002.{i;, 0.022 0.0/8 0.0/4 0.0/0 0. OOrtJ

PAGE 78

76 Curve II of' Figure 5 demonstrates that there is a maximum oxidation rate tor the reaction mixture that is very sharply def"ined 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 of"t quite steeply on both sides. Toward the acid side the constant .attains a considerably lower value than on the al kaline side; . and 1t. appears to be destined f"or a value ot ' ., , '' , aero at a pH someiirhere around a.5. This is to be expected ,' since Curve I indicates that the oxidation goes completely ,. . ., I ' to sulf'a te within . the f'irst minute or reaction time 1ft a pH , ' value in the general vicinity of 4.0. On the alkaline side . . . . . . or 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 ~he value or constant l!, f'or the experiment conducted at pH,7.92 has been disregarded in draVing Curve II. Table 16 discloses that the solutions used in this particular . . . experiment were calculated to have ionic strength values ot 0.2154 as compared to values ot 0.1093 tor the solutions 'employ~d in-the other experiments of' this study •. This 'dif"terence in ionic strength values could not be avoided in ~e investigative methods used, and later work concern-ing the effect of' ionJ.c strength on the reaction justif"ies the disregarding or the above-mentioned value. In Figure 6 some of' the experimental data trom Tables 9

PAGE 79

-7'1 FIGURE. 6 EFFECT OF HYDROGENJON CON CENTRA TJON ON THE RATIO OF CHLORINE ro SllLF/O REACT 36 J.4 .3.2 '-..: 3.0 2.8 -~ ' c.ri, ';.__ () ?..4ti ..... 2,2 () 2.0 /.8 /.fiJ ( See Tables 9 f 14) Curve I: Reachflll lime is I min. Curve .a: Raacfion hrne is /0 min. Curve m: I/me i.s 40 min. 5.0 14.0 70 pl! Ill. JI I 8.0 ~-0

PAGE 80

78 and 14 are presented in a way designed to illustrate with greater clarity some or the existing relationships that are not read1ly apparent .from aneumination or Figure 5. The experimental ratios obtained with reaction times or 1, 10 and 40 minutes, respectively, have been plotted against pH values to give an interesting family or curves. Curve I or Figure 6 and Curve I or Pigure 5 may be regarded as being the same, each or them showing_how the experimental ratio tor a one minute reaction time varies with pH. However, in the case or Figure 5 the curve was derived from calculated data, whereas in Pigure 6 all of the curves represent the observed data. Thus what has been said about Curve I or Figure 5 may be applied to Curve I or Figure 6. In tact, 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 or 10 and 40 minutes, respectively. It is observed that these latter curves are displaced upward in relation to CUrve I, wb.1.oh is to be expected when the longer reaction times are con sidered. The most striking feature or the two curves, how• ever# is the bumped area in each or them in the vicinity or pH 7.0. The depth or the bumped area is seen to be a f'unotion or the reaction time, which further illustrates the maximal oxidation rate in this pH range. A significant feature or this maximum in the oxidation rate is readily recognized from the relationships that exist between the curves presented

PAGE 81

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 or 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 2.6 is indicated tor a one minute reaction time at a pH or about 5.9., but this same ratio can be attained at a pH value or about 7.0 by allo'lling a reaction time or 40 minutes. Vere it not tor the max:Jmal oxidation rate in this pH range it would re quire a muoh greater reaction time to attain the given ratio at the pH value or 7.0. The manner in which the experimental ratio increases with the hydrogen-ion ooneentration, as illustrated in Figure a. is suggestive ot the way in which the concentration ot undissociated hypoohlorous acid must increase with hydrogen ion concentration in a solution or chlorine water. It an analogy between these two relations could be drawn it might give some indication as to vb.ether or not tree hypoehlorous acid is the effective oxidising agent in the reaction. Consequently., calculations have been made according to the .. method su,ggested by McKinney (53) to illustrate the relation ships that exist between the activity f'ractions of' free hypochlorous acid and the h.fpochlorite ion and the hydrogen ion concentration. The method that is involved is as follows:

PAGE 82

... 80 For hypoohlorous acid uc10 nl I c10aglxac19~ = 5.6 X 108 8 BC10 Let aHClO = l . aiJ : _ :1o~PH 0 IIC10: aClO .. = 1 : : l t 5, I l!r 8 10-PH : 1: 5.6 x lO(pB;. 8)' Let aBClO f ~c10 : aT . anc1ofaT.: 1 /. 5 6 x 1 10 (pH -s) = 1,' 10 (!u 7.252) ac 10 -/aT : . ,6 x 1q
PAGE 83

pH 3.2 3.4 3.6 3.8 4.0 4.2 4.4 4.6 8 5.0 . 5.2 5.4 5.6 5.8 6.0 6.2 6.4 6.6 6.8 7.0 7.2 7.4 7.6 7.8 s.o 8.2 8.4 s.s 8.8 9.0 9.2 9.4 9.6 9.8. 10.0 10.2 10.4 10.6 10.s 11.0 11.2 1 I lJr,4 TABLE 17 ACTIVIff FRACTIONS HYPOCHLOROUS ACID 1.0000 0.9999 o.9998 0.9996 o.9994 0.9991 o.9986 0.9978 o.9965 0.9944 0.9912 o.9861 0.9782 o.9650 o.9470 o.9185 0.8764 o.8177 o.7391 o.6410 o.5299 o.4149 0.3096 0.2208 0.1515 0.1013 . 0.,0662 0.0429 0.0276 0.0115 0.0112 0.0070 0.0045 0.002s 0.001s 0.0011 0.0007 0.0005 o.oooa 0.0002 0.0001 , .. o.oqog 81 0.0000 0.0001 0.0002 0.0004 0.0006 0.0009 0.0014 0.0022 0.0035 0~0056. o.ooes 0.0139 0.021s 0.0341 0.0530 o.os15 0.1236 0.1823 0.2609 0.3590 0.4701 o.5851 .6904 0.7792 0.8485 . 0.8987 o.aaas 0.9571. o.9724 0.9825 o.esss o.9930 0.9955 . o.9972. o.tss2 0.,9989 0.9993 o.9995 . 0.9997 0.9998 0.9999 ,1,,ooqo

PAGE 84

82 FIGURE 7 ACTIVITY FRACTIONSHYPOCHLOROl/S ACID (See Tobie 17) Curve I: Achvdy frocNon of und/ssociafed hypoch/orous _ dc/d. Curve II: Aclivily froclion of 1rpochlorlle ion. J. I.Or-----0.8 -1:: ,'\:) 0.6 } ' ':I,. \l o.+ (}, C. JI s.o JI

PAGE 85

83 this Ticinit7 the aetiTity traction or the hypoehlorous acid and the experimental ratio both increase sharply, and the curTes are quite similar in appearance 1n the alkaline region., the similarity being particularly striking in the case or CurTe III of' Pigure 6. However, in the more acid ran&e the curves bear little resemblance to one another. The activity fraction curve or hypochlorous aeid begins to taper off' at pH 6.5 and to approach a ma::id.mum value ot unity; whereas the experimental ratios at this pH value are incr,~asing even more sharplywith decreased pH •. ltteai of Iwc lirepgth on the Reaction The previous observations eoncernirg the ettect ot hydrogen-ion concentration on the reaction serve to emphasise the necessity for rather severe pH control throughout the various experiments included in the investigation or this . oxidation. It has been mentioned in connection td.th the preparation of' the concentrated buf'f'er 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 o.05 pH unit, and it was f'ound that conaiderable concentrations of' buffer materials were required 1n the work ing solution~ to limit the pH change to such narrow limits. Consequently, it was necessary to study the oxidation 1n butf'er solutions having relatively large ionic strength values. It has been seen in the preceding work that an ionic

PAGE 86

84 strength value.of' 0.1093 vasadopted,as a convenient and constant value to be used in most of' those experiments where the ttseot." solutions haring such.an ionic strength.wall pos sible. However, inf"ormation was desired relative to the . reaetion:1n soluti•s having ionic strength values comparable ) to, those values,f'ound in natural waters, such,values bei~g in the. general vicinity of'.005. In order that. such intorma,tioncould.be obtained,. experiments designed. to ' . . . . ' . ' indicate the ef'f'oct of' ionic strength on the .reaction. vere , J' eondueted over a nde range of' ionic strength values with the thought that the results could be extrapolated to leaser ionic strength values. The procedure f'or these experiments .has been described. earlier. It .has been pointed out 1n the description of' the pro cedure.f'or the.experiment~ .dealing With the ef"f'ect of' ionic strength on the reaction that the investigative methods. used were similar to those used in determining the etf'ect of' ,, . , . . . . . time. The notable ditf'erenee was that a means vas provided to vary the ionic strengths ot the working solutions between the 1ndirtdual experiments while leaving the pH value ~onstant. All.of the _experiments were conducted in solutions having a pH value that was as nearly constant as possible .at 7.0. This Yalue was choaen because it lies at the middle.of' the range of' pH values considered in this inveatigation, and ~t l , 1s possible to.attain a lower ionic strength value 1n :solutions at this pH value than others With the assurance

PAGE 87

85 that the pH change due to the addition or reagents rill be kept within the prescribed minimum or 0.05 unit. Further,. _, more, it can be seen from Figure 5 that the selected value ' . ' l a .,_ also liea within a range of' pH values where the e.rtect due to small variations in the hydrogen-ion concentration is ' ' . ' practically negligible~ . The results observed. during the ioni~ strengt~ expert... ments are shown 1n Table 18. It is note4 that the pH value > ., ' •! ,. decreased steadily.from 7.05 to 6.89 as the ionic strength was increased f'rom 0.05483 to o.2046 during the ~xpe_riments. However. the previously accomplished work concerning the ' ' _ef'rect or.h1drogen•ion eoncentration indicates that this , , l small variation or 0.18 pH unit produces no inaocu.racies dm ., . ' ', ... , . ' . ' ' . to that etteot in determining the errect or ionic strength on the reaction as long as the experiments ar~ conducted within the range or pH values lying between 6_.5 to 7 .3. The constants, .!! and h, of' the empirical equation,, ' . R atb, were calculated ror e~oh or the experimental. runs in the study# and the resul. ts are reported in Table 19. An . examination or the results indicates that the initial in\ crease in ionic strength apparently produced a marked in. crease 1n the Yalue or constant L but t"ur~r increases in the ionic strength value had no recognizable erre~t on.this constant. A glance at Table 18 shows that the initial in crease in ionic strength was accompanied by a change in the chloride concentration or the working solutions f"rom Oto

PAGE 88

Ionie Strength 0.05463 1f 0.07963 .. " ti 0.1046 TABLE 18 EPPECT OF IONIC STRENGTH ON THE RATIO OF CHLORINE TO SULPIDE REACTED. OBSERVED DATA. Volume ot butter solutions Sultide addedt Concentration or sultide solution: Chlorine addodi Temperature: Reaction Time
PAGE 89

TABLE 18 Continued Ionic Reaction pH Chloride Chlorine Chlorine to Chlorine to Strength Time Concen ira tion Reacted Sulfide ' Sulfide (mtnl C1vm2. <•> . B1t12 'by&) . Ratiq Cmoltl, , 0.1296 1 ;6.95 2640 10.09 5.-05 2.29 5 10.20 5.10 2.31 10 10.95 5.48 2.48 20 11.13 5.57 2.52 40 tt 11;32 5.66 2.56, 80 ' .. ' . ~64-. 5.82 2.64' 0.1546 1 6.93 3520 ~.96 4~98 2.26 I 5 10.47 5.24 2.37 o:> ,. 10 1().83 5.42 2.46 ,_. 20 11.13 5~57 -2.52 t 40 ll.15 5~58 2.53 80 11~48 5.74 2.60 ' 6~91 10~09 5~05 0.1796 1 4400 2.29 5 10.34 6.17 2.34 10 10.12 5.36 2.43 ,. 20 10.90 5.45 2.47 40 10.94 5.47 2.48 80 11.13 5~57 2.52 0.2046 1 6.89 5280 10~20 f>.10 2.31 5 , . 10.16 5.08 2.30 10 ,. 10.73 5.37 2.43 " 20 " 10.ss 5.43 2.48 40 ,. .09 5.55 2.51 ft 80 '11129 . 51 21

PAGE 90

Ionic Strength 0.05463 " 0.01003 rABLE 19 Eff.BCT OF lONIC STRENGTH ON . TBE RATIO OF CBLOIUNE TO SULPmB REACTED, OALOULATED DATA. Souroe of data tor-caloulationsi table 18. Basis for calculations1 R: at, where Reaction Time (min) 1 5 10 20 40 80 .l . 5 10 20 40 80 n : Chlorine to sulfide ratio (mole)• . . t = Reaction time (min), . . a.am l?. are oharaoteristio constants. Conetant A 2.04 ft " .2,30 .. Constant ll. 0.0497 It .. ... 0.0526 Calculated Chlorine to Sulfide Ratio (mole) 2.04 2.21 2,29 2,37 2.45 2.54 2.ao 2.51 2,60 2,70 2.so 2,90 1' Deviation , , trom Observed Ratio I co co I ..;.7.5 1~4 -1.3 o.o -o.4 o.o -0.9 _ .. .5 o.o 0.1 o.4 o.o

PAGE 91

TABLE 19 Continued Ionic Reaction ,, Constant Constant Calculated i Deviation Strength Time a l?. Chlorine to trom Observed (min) Sulf'ide Ratio Ratio (mole) 0.104a 1 2.23 o.0492 2.'23 -3.9 5 2.42 2.5 10 2.50 -o.4 20 2.59 1.9 tt 40 .. 2.68 o.o . ' 80 2.77 ., o.o tl) co 0.1296 l 2.2s 0.0331 2.28 -o.4 . I 5 2.41 4.2 ft 10 tt 1h47 -o.4 ff 20 2.52 o.o 40 ft " t.58 p.s 80 l.64 . o.o o.1546 l 2.27 0.0307 ;a.27 0.4 It 5 2.39 o.s II 10 ff ft 2.44 .. (l.8 ft 20 1t )?'.49 -1,.2 40 2.55 o.s 80 " ., ft.60 o.o

PAGE 92

TABLE 19 Cont:lnued Ionic Reaction. Constant Constant Calculated 'lo Deviation Strength Time !. 12. Chlorine to f'rom Observed (min) Sulfide Ratio Ratio (mole) 0.1796 1 2.28 o.02as .28 -0.4 5 " 2.37 1.3 " 10 " .41 -o.s l 20 .45 -o.s (0 40 tt 2.49 o.4 C 80 2.53 ... .0.4 I ...... ,i, .0.2046 .. 1 2.29 0.0251 2.29 -0.9 5 ft " 2,.39 .s tt 10 ff ft 2.43 o.o 20 " tt 2.47 0.4 40 " ft 2.52 0.4 80 It ft 2.56 o.o

PAGE 93

91 880 parts per million. Literature previously cited (42) suggests tie possibility that the observed e:r:reot on constant A may be due to this addition or chloride, and this point was investigated further by later experiments planned to demon strate the ettect o:r chlorides on the reaction. Constant b was observed to exhibit a de:finite variation with changing ionic strength. The relationships that exist between the two are represented graphically in Figure a. It appears that J?. has a constant value in the vicinity or 0.000. within the limits o:r experimental error, :tor ionic strength values 0.05463, 0.07963, 0.1046 and the prenously studied value or 0.1093. However, a sharp decrease oc~ur• in the value or the constant with increasing ionic strength values just in excess or this range, until it appears that a minimum value or about 0.022 to 0.023 is approached at an ionic strength value somewhat greater than 0.20. An extrapolation or the data shown in Figures to the range or ionic strengths normally encountered 1n natural waters suggests that the re action is tmaf'f'ected by variations in ionie strength in this region. In tact, it is indicated that as tar as the ef'tect or ionic strength is concerned, the results or most of' the preceding experiments are applicable to such dilute solutions as those f'ound in natural waters. lt is pleasing to note that this study explained the apparent discrepanc7 in the value or the constant]!; tor pH value 7.92 in Curve II of' Figure 5. All of' the other points on this curve were de

PAGE 94

92 Fl6UR 8 EFFECT OF IDA;;~STRENGTH ON THE CONSTANTS OF ' THE EMPIRICAL . E9llATIONj R = 1 afh_ (See Tohle 18) 0.07/J 00'5 00.55 0 0.050 0015 ~I ' 0.040 "\..; If) 0.015 P.030 ~020.__---'---'----'---'-----'--'-----'L----l---L----l----.JL---....J o 0.02 00-1 ao6 o.oa i 0.10 0.1~ 0.14 1.10 0.1a 0.20 022 Ionic Sfrengfh

PAGE 95

93 termined in solutions having a common ionic strength value or 0.1093, but this particular point in question was de termined at a value or o.2154. It is seen from Figure 8 that decreaaing the ionic strength trom 0.21 to 0.11 has the ef'tect ot increasing the value or constant l!. by a tactor of' 2.2. Applying this taotor to the point in queation in Figure 5, it is seen that the value obtained tor the constant is brought into very good agreement with the pre supposed value. In view or theroregoing evidence that the ettective agent in this oxidation reaction is hypoohlorous acid, it was thought adviaable to investigate the etteot or the con centration ot chloride ion, which was used to vary the ionic strength values during these experiments, on the hydro lysis or chlorine to hypochlorous acid. The equation tor this hydrolysis reaction is as rollovs: c1/. ao10; R; 01 2 I u 2 o Lat.imer (29) gives -6,315 calories per mole as the f"ree energy change tor this reaction, which yields an equilibrium constant of 4.27 x 10 4 Using this value tor the equilibri.um eonetant, it is calculated that at a pH value ot 7.0 and With no chlor.ide added all or the oblorine that vas added to the working solutions involYed in the experiments existed in the solutions as hypochlorous acid. This acid was ionized in accordance with the relationships shown in Figure 7. Further calculations show that when the chloride concentration in the

PAGE 96

working solutions was 5,280 parts per million.,. the largest eonoentrationused in these experiments., the total.concen tration of' hypochlorous acid was decreased by only 0.2~ •.. It is assumed, therefore I that t,he effects. noted in this study are not due to any shift in the chlorine hypochlorous acid equilibrium caused by the addition of' chloride ion. ltt,ct o[ jllilor&de Cgpeentrati2n on the Reaet&9n Potassium chloride vas added to the buf'f'er solutions employed in the preceding experiments as a means f'or 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 1nwater are very nearly the same. Furthe~ore, the ions. are affected neither by the oxidit:ing.action of' such strong oxidizing agents as chlorine nor the reducing action.of such reducing agents as sulfides. trowever, certain of' the results from the ionic strength experiments, which have been pointed out., su,ggest the strong possibility that ehlorides affeot the oxidationreaetion in some manner, and it has already been mentioned.that Higgins (42) noticed a.stimulating et.feet produced by chlorides on the bleaching action or chlorine 9olutions •. Consequently, experiments were planned with which to study this possible effect on the reaction., The procedure for these experiments has been described in the •eotion concerning the experimental methods, and again at

PAGE 97

95 tention is called to the similarity between these experiments and those used to determine the etteet or time. The pH and ionic strength values vere held constant throughout ~he . experiments at values ot.0 and 0.1093, respectively, in order that the resultsof' the experiments voutdbe a compar able.basis with the results or preceding ones. The observed data f'rom the experiments concerning the effect of' chloride concentration on the reaction are pre sented ui Table 20, and the calculated data are shown in Table 21. The calculated data relative to eonstants .a and!!, of' the empirical equation, R = atb, are briefly summarized f'or convenience as f'ollows: Chloride Constant Constant % Deviation or Concentration a l!. J! ~om mean (p:em) 0 2.01 0.0477 -5.2 385 2.07 0.002a 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 2,9&1 3.6 Average 0.0503 Inspection of this summary shows that the :tull ettect 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 ot constant.!. with the add1.tion ot chloride•~ whereas the

PAGE 98

Chloride Concentration .. {pem) 0 385 .. TABLEao EFFECT OF CHLORIDE CONCENTRATION ON THE RATIO OF CHLORINE TO SULFIDE REACTED., . OBSERVED DATA •. Volume ot buf'ter Sulfide added: Ooncontration of Chlorine addedt Temperature:, Ionic strength: Reaction Time Cm1n> 1 5 , 10 20 40 80 1 5 10 20 40 80 solution:: sulfide solution:. 2 mg/529 m1 or 525 ml~, 2.oomg •. 3.78 ppm. 17.68mg., 25o c •. 0.1093 pH. 6.97 ff fl " lt 1t Chlorine Reacted (mp:) 9.17 9.58 9.58 10.22 10.45 11.07 9.27 9.98 10.16 10.64 10.89 11.so Continued Chlorine t.o Sulfide Ratio ,m,l 4.59 4,.79 4.79 5,.11, 5.23 5.54 4.64 4.99 5.08 5.32 5.45 5.90' Chlorine to Sultid& l Rat:{9' (190;\e):_ '~ ' 0) a.os I 2.17 2.-17 2.32 2.37 2.sr 2.10 2.26 2.30. . 2,.41" 2.47 2.67

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TABLE 20 Continued Chloride Reaction pH Chlorine Chlorine to Chlorine to Concentration Time Reacted Sul.fide Sulfide (ppm)
PAGE 100

Chloride Concentration (ppm). 0 " tt 385 " " " .TABLE 21 ' EFFECT OP'CBLORIDE CONCENTRATION ON THE RATIO OF CHLORINE'. TO SULPIDE REACtED. CALCULATED DATA. Source'ot data tor calculations: Table 20. Basis tor calculations~ . R atb, Vhere B.: Chlorine to ault~deratio (mole). !, Reaction time (~n), ' A a~ J?. are characteristic constants •. Reaction Constant Constant Calculated Time Jl l!. Chlorine to (min) sulfide Ratio (moJ.e) 1 2.01 ,0.0477 2.01 5 2.17 10 It " 2.24 20 ., 2.32 40 It 2;40 80 .. 2.48~ l 2.07 0.0523 2.07 5 2.25 10 If H 2.34. 20 2.42 40 2.51 80 2.61, .. Continued 'I, Deviation f"rom Observed Ratio I CCI 00 -3.4 I o.o 3.2 o.o 1.3 -1.2 -1.4 ().4 1.7 0.4 1.e . -2.3

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TABLE 21 Continued Chloride Reaotion Constant constant Calculated % Deviation Concentration Time a li Chlorine to tromObserved (min) Sulfide Ratio 0 II Ba tis I (uJ.e > 769 1 2.07 o.0464 2.01 -0~5 .. 5 It '2.23 -1.3 t, 10 tt " 2.30 2.2 It 20 .. .37 -1.7. ft 40 1t 2.45 o.o " 80 " .. 2.53 1.2 1150 1 2.07 . 0.0533 2.07 -1.0 l ft 5 " 2.25 -o.4 ff 10 ff 2.34 1.7 U) ff 20 " 2.42, , .. 3~4 , tt 40 " fl 2~52 -o~s tt 80 " 2~61 -2~6 1540 1 2.16 0.0502 2~16 o~o .. 5 ft ,. 2.34 -2~9 ff 10 It a.43 2~9 ft 20 tt It 2.51 ,. o.s 1t 40 " 2.60 -1;1 80 " ft 2~69 o~o 1920 1 2.17 o.os21 2~17 -0~5 ft 5 If 1t 2.35 1~3 It 10 .. ff 2:44 -o~s 20 ft 2.53 o.o ft 40 ft 2.63 .:o.4 80 tt 1,7! 0,2 I 7

PAGE 102

100 value or constant h remained constant. Within the limits or the experimental error. It is indicated by these results that the increase in the value or constant.!. as the chloride concentration increases is stepwise. Bovever. in view or the relatively small overall change 1n the Talue or the constant and the experimental error involved, it cannot be stated det1nitely that this observation is significant. The results or these experiments otter still more proor that the variations or constant h with ionic strength observed in the preceding ionic strength experiments were not due to the increasing chloride concentration in the working solutions. They also show that the initial variation or cons~ant !!. that was noted 1n those experiments was caused by the addition or chlorides. It is gratifying to see that the addition or potassium chloride to the working solutions or preceding . . experiments caused no change at all in the values or constant l?. and but relatively little change in those valuu calculated

PAGE 103

DISCUSSION OF RESULTS The titrations or sultide solutions with chlorine water that were made during this investigation otter evidence that the oxidation ot sulfide with chlorine proceeds first ot all to the tree sult'ur state, vhich bears out the general con tention of' other .investigators that sulfur is the primar7 product of the reaction. However, .it was seen f'rom the experiments concerning the etf'ect of' concentrations on the oxidation that the ratio of' moles or chlorine to sulfide re acted is always greater than unit7, even when there is an excess of sultide in the solution. Th.us it was shown that when the chlorine dosage is added quickly to the sulfide solution by the teehnic described tor this in-vestigation a '' part or the primary product must be_ oxidized simultaneously to a hj_gher oxidation state by some secondary reaction. Since it has been ahovn that the only end-products or the reaction are f'ree sulfur and sulfate (46), the final product or this secondary reaction must be sultate. The results or certain or the experiments that constitute this imestigation indicate that the initial, sulfur producing reactions may be regarded as being instantaneous in action# and these reactions are postulated to be as rollows: B2S / HClO : S / u/. /. 01/. H20 usI c10= s I c1/. ou101

PAGE 104

102 The first or these reactions takes place in acid solutions with a standard tree energy change of -62#500 calories, and the latter one occurs in alkaline solutions with a corre sponding change of -85,500 calories. Bullock and Porbes (16) claim on the basis of experiment that the sulfur formed by the above type or reaction is set rree in an active state that is quite susceptible to further oxidation. With this point in mind, Choppin and Faulkenberry (46) suggested that the active sul:tur formed by the reaction between sulfide solutions and hypochlorites may react in two ways; it may combine With its~lr to become relatively inactive diatomic sulfur by the reaction, S /. S : s 2 or it may be oxidized to sulfate by hypochlorite in a series or reactions which they presented. Their reactions appear to explain successtuily from a kinetic point or view certain or their experimental data, and this latter oxidation is re garded as being the secondary, sulfate-forming reaction which has been found to accompany the initial oxidation of the sulfide. However, the probability was evidenced during the present investigation that f'ree bypochlorous acid is the active agent in this secondary oxidation process rather than the previously suggested hypochlorite ion. Therefore, the liberty or making this change in the reactions presented by Choppin and Faulkenberry has been take~. Thus the overall oxidation may be expressed by the reaction,

PAGE 105

103 s /. aHc10 I u 2 o = so 4 -/. 5Bf/. 301and the meehant•m or the process i . s explained .. bl' the .t"ollov1ng reacitions: s /. uc10 = so I al /ci so / H20 : HaB02 u 2 so 2 ,J. uc10 : u 2 so 3 '; al j c1u2so3 ,J. nc10 : so 4 -I a-J /. c1These reactions still of'f"er suceessf'llt . explanat.:lons tor the data ot Choppin and Faulkenberry" and they are at the same time more eonaiatent Vi.th regard to the results or the present investigation. A ratio or tour moles or chlorine per mole or aulf"ide was 119ed in the large majority or . the experiment• employed to study the ef'.tects of" the ditrerent variables on the oxidat:Lon of' 81.tlf'ide bf chlorine du.ring the course ot this inveetigation. This ratio us round to yield an excess or chlorine in the 11olution throughout all the reaction times employed in the experimental vork. In view or this t'act and the f'indinga that the reaction involving the oxidation o:r aulf"ide to sulfur takes place immediately with the aecompani ment ot the aecondar7 reaction whereby a portion of' the sutt'ur 1• oxidised to aulf'ate, it is aeen tha't the reaotion mix .. turea must have contained sulf'ur, sulf'ate and chlorine at the in•tant af'ter the chlorine was added and mixed vith the solution.., Since the ertecta of' the di.ff"erent variables that were noted during the investigation were f'or the most

PAGE 106

104 part erreets on the system atsome time af'ter.tbe mmediate reactions had taken place, it is understandable that the major portionorthe experimental work conducted during these.studies must in reality have concerned the secondary oxidation-process in which the reaction wasbetveen sulfur nnd chlorine. The .fact that this secondary reactfon was seen to progress with time in those experiments where there was an excess-or chlorine leads to the conclusion that it is a consecutive reaction in the presence or excess chlorine, as well as a concurrent one# in relation tothe initial oxidation. The results o.f the investigation show that rrom the standpoint or water works practice the hydrogen-ion con centration is by ~ar the most important variable studied, since it exerts such a considerable influence on the dosage or chlorine that is required to insure comple'teness or the initial oxidation and• there.fore~ complete elimination or the sulride. For example. td.th the conditions employed in the concentration experiments it.was seen that a dosage of' about rour milligrams or chlorine per milligram or sultide was required to produce a slight exoess ot chlorine in the reaction mixture at a pH value of 9, whereas the dosage re quired to yield an excess or chlorine at a pH:value or 5 was round.to be around 8 milligrams ot chlorine per milli gram of' sulf'ide. This variation with hydrogen-ion con centration or the chlorine dosage required to eliminate

PAGE 107

... 105 sulfide in the reaction mixture actnally retlects the erteet or bydrogen-ion.oonoentration on the simultaneo~ secondary oxidation process ,that produces sulf'nte. Thus it is seen that.although lover pH values require higher chlorine dosnges, they yield correspondingly larger sultate to,sultur ratios in 1 . . . the reaction mixture. . . These relationships may be of' consider able userulness.~ water works practice where it is advisable . ' '. : ' ' : '. , ' . '-' . , , to strike a l>alanee .between the. cost ot lowering the pH, the ' : ' ' . ' . . ,. . cost ot chlorination and the amount of' f'ree sulf'ur that ~s allowable in the finished vater. In this latter .c~nnection . . . . . '. t. ; . it .was .noticed in those experiments of' this investigation •, ., ' ' l 1• where original sultide concentrations of' 3.78 parts per million and chlorine dosages ot.four moles or chlorine per ,, . ' '' . ' ; 1 ' , , ' ,.. . ' .-, mole or.sulfide were utilized that free sulf'ur was.visible by casual examination throu,ghout all except the experiments in. which.the pH value is around 5.o. It is,not likely that these particular relations hold ror other concentrations or sulfide •. However~ it is estimated f"rom the above observations and the data 1n Tables 9 and 2 that the threshhold or visibility or rree sulf'ur in water by a casual examination or a.550 milliliter sample lies between 0.8 and 1.5 parts per million. Next in importance from the standpoint or water works operation appears to be the eftect or reaction time. Although the initial reaction wherebythe sulf'ide is actually eliminated takes place immediately when suf"fieient chlorine is added, the ,. ' '' secondary oxidation or sulfur has been round to . progress vi th

PAGE 108

106 time when there is a slight excess of' chlorine in the re action mixture. As seen in Figure 6, the etf'ect is not very considerable except in the pH range 6.0 7.4, within which region a maximum in the reaction rate occurs. It has been previously pointed out that the increased reaction rate in this pH range makes it poasible to accomplish conveniently at a higher pH value, by allowing a longer reaction time, the same degree or oxidation in the reaction mixture that can be accomplished at a much lover pH and shorter reaction time. The optimum pH range in which to take advantage of' these relations is seen f'rom Figure 5 to lie between pH 6.5 and 7.3, where the reaction rate is actually at its maximum. It is judged f'rom the results of' this investigation that the etf'ects of' temperature and 1onJ.c strength on the reaction are of' little or no consequence in the range or values normally f'ound 1n natural waters. The presence of' chlorides was seen to increase the extent or the reaction-that takes place within the rirst minute# but the experiments concern ing this ettect ottered inconclusive results in regard to the exact nature ot the increase in the range or chloride concentrations f'ound in fresh water supplies. The increase is relatively small, however# and is probably of' no real signUicance f'rom the standpoint ot water works practice. Furthermore, it is included in the experimentally determined chlorine demand or a sulfide bearing water. It is not in tended that the results or this investigation vill make it

PAGE 109

107 possible to eliminate th• chlorine demand test tor determin ing the chlorine dosage required for a sulfide bearing water# since the water Will usually have a demand tor chlorine that arises rrom the presence or other substances. Hence each water oonstituies an individual problem. However, it is hoped that the results or these studies will complement chlorine demand tests and sulf'ide determinations in arriving at the most satistactory way to treat such waters.

PAGE 110

StJMKARY A series ot studies was conducted to determine the etteots ot'a number ot variables on the oxidation ot Yery dilute, aqueous, sulfide solutions by chlorine. All the concentrations ot sulfide employed in the studies were in the vicinity of' 2 6 parts per million. The overall re.;. action, which is independent or sulfide concentration in this very dilute region, involves an initial oxidation or the sultide to sulfur and a secondary oxidation ot a portion ot the tree sulfur to a higher oxidation state that has been shown to be sulfate by previous investigators. The initiel reaction takes place immediately with the addition or chlorine to sulfide solutions. 'When the chlorine is added quickly the secondary reaction takes place simultaneously with the initial one. and with an excess or chlorine in the reaction mixture it continues slowly as a consecutive reaction. Thus the chlorine dosage required to completely eliminate sulfide lies between the stoichiometric quantity calculated to pro duce sulfur and that calculated to produce sulfate. This dosage is markedly dependent upon the hydrogen-ion concen tration of the solution. The extent of the overall oxidation is dependent upon the ratio or chlorine added to the sulfide as well as upon the hydrogen-ion concentration of the solution. It increases 108

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109 substantially with increased chlorine to sultide ratios in acid solution;but this increase becomes less marked as the pH is raised and appears to become stabilized at a pH value around s. Vb.en a ratio or• f"our moles or chlorine per mole or sultide(S.84 milligrams per milligram) is employed the reaction progresses oonsiderably•beyond the.sulfur stage dur ing the :first minute of' reaction, •the extent depending upon the hydrogen-ion concentration. For reaction times between. one and eighty•minutes the reaction proceeds in accordance with the empirical equation, R: atb; where!! is the ratio otohlorineto sulride reacted in moles, .1s the reaction time in minutes and .a. and J?. are characteristic constants. Constant 1!. reelects the extent or the reaction ror a reaction time or one minute, and constant.Ji reflects the reaction rate. A study made to determine the effect of hydrogen-ion concen tration on these constants throughout the range of pH values :from 5 to 9 shows that constant.!. decreases steadily from a value or 3.4 at pH 4.9 to an apparent minimum value or 1.7 at about pH9.0. In the case or constant l!, a broad maximum was discovered f'rom pH 6.5 to pH 7.3, which indicates a maximum reaction rate tor the secondary oxidation between these pH values. The values for the constant decreases . aharplt on either side orthe maximum, and it appears that the constant approaches a minimum value at about pH 9. The conspicuous similarity between the manners in vhieh the extent or the reaction and the activity ~action or free

PAGE 112

--110 hypochlorous acid vary nth hydrogen-ion concentration points to theprobability that rree hypochlorous acid is the active agent in the oxidation~ Temperature variations in the range between 15 and 250 c. have no signtticant er:tect upon the oxidation. Variations 1n ionic strength have noapparent ettect up to vnlues or about 0.11. However, the reaction rate or the secondary oxi4ation decreases sharply as the ionic strength is increased rrom o.11~ until a minimum ~s apparently approached at a value around 0.22. The presence or chlorides 1n the sultide solutio~ causes a small but noticeable increase 1n the extent or the immediate~ overall oxidation, but the rate or the secondarJ reaction arter a ono minute reaction time 1$ unaff'ected. Attempts to determine the etteet or varying the chloride conoentration led to inconclusive results. Prom the standpoint of' water works practice the hydrogen ion concentration is by :far the most imp9rtant :factor to be considered 1n the reaction between sulf'ide solutions and chlorine. The ef'f'ect or reaction time is worthy or con sideration, but the ef'f'eots of' temperature., ionic strength and chlorides, as they are commonly encol.llltered in natural '.m.ters., are rather insignif'icant in character.

PAGE 113

BIBLIOGRAPHY (1) R. L. Derby, J. Am. Water Yorks Assoc., 20, 813., (1928) (2) c. R. Cox, J. Am. Water Yorks Assoc., 28., 1855, (1936) (3) s. Rideal', Trans. Faraday Soc ., i, 179 (4) C. P. Hoover, •water Supply and Treatment•, 5th Edition, National Lime Association,. Washington., D. c.', (1943),. p. 113 (5) R. Pomeroy and F. n. Bowlus., Sewage Works J., is, 597, (1946) (6) s. T. Powell and L. G. von Lossberg., J. Am. Yater Works Assoc.,~ 1277., (1948) (7) "Standard Methods f'or the Examination of" Water and sewage•, 9th Edition, American Public Health Association, Nev York, (1948), p. 103 (8) Wallace and Tiernan Co., Inc ., Brit. Patent No. 551,645 (1943) (9) H.B. Dunnicliff" and c. L. Soni,, Proc. 15th Indian Sci. Congr., 1928, 167 (10) H. B. Dunniclif'.t and c. L. Soni, J~ Phys. Chem., l!., 81., (1929) (ll) H.B. Dunniclif".f, G. S. Kotwani and M.A. Hamid, J. Phys. Chem., 39, 1217, (1935) (12) M.A. Hamid, G. Singh and H.B. Dunniolitt, J. Indian Chem. Soc., l&, 595, (1935) 111

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112 (13) H. B~ Dunniclif'f' and B. Prakash, J. Indian Chem. Sooo, J,l, : 505 (1~35) (14) ~ ~. Hamid, v._s. Bhatia and H. B. Dunniclitr., J. Indi~n Chem. S~e ., ll, 697., (1936) . . . (15) H. 1:'• Dunni~lif'f',, J. Indian Cllem. Soc., Ind.&. News Ed., 2., 45 (1944) (16) J. L. Bullock and G. S~ Forbes., J. Am. Chem. Soc., 55, 232 (1933) . ' (i7) I.M. Kolthoff" and E. B. Sandell., •Textbook of' Quanti ~t:f:ve Inorgani.c Analysis",. The MacHillan.Company,_New York,. (19'4:7), p. 634 (18) H.F. Prost, Analyst., !12., 90, (1944) (19) M. Pollaeci., Mon. Soi., _(4), &&, 373 (20) s. A. Sh~huknrev. and E. M. Kireeva-Tuzulakhova, J. Gen. Chem. (u.s.s.n. ), L No. s -s., 1125, (1931) (21) L. s. Bagster, J. Chem. Soc ., J.928, 2631 (22) s. Mohammad and G. s. Ahluwalia, J. Indian Chem~ Soc ., ll, 39$} .,_ ( 1941) (23) S. Mohammad ands. N. Bedi, J. Inclian Chem. Soe., _&l, 55, (1944) (24) H. B~ 'Dunniclif'f' and G. s. Kotvani, J •. Phys. Chem 35, 3214,. (1931) (25) R~ S~ Denn,.J •. Am. Chem. Soc., 40, 619, (1918) (26) H. D. Dunniclif'f' ands. Mohammad., J. Phys. Chem.,. U, 1343 (1929) (27) R. K. Bahl ands. Singh, J. Indian Chem. Soc., li, 339, (1939)

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113 (28) H.B. Dunnielirr ands. D. Nijhawan, J. Chem. Soc., 1926, 1 (29) w. M. Latimer, "The Oxidation States or the Elements and Their Potentials in Aqueous Solutions", Prentiee Hall, Inc. (1938). (30) I. M. Kolthoff and E. B. Sandell, •Text book or Quanti tative Inorganic Analysis•, The MacMillan Company, Nev York, (1947), p. 614 (31) M. Kapp, Bull. soc. encour. ind. ntl., 131, 330, (1932) (32) H. v. Tartar, J. Am. Chem. Soc., 12., 1741, (1913) (33) H. Bassett and R. G. Durrant, J. Chem. Soc., 1927, 1401 (34) A. Skrabal, z. Anal. Chem.,.!, 107, (1924) (35) s. H. Higgins, Chem. Soc. Proc., &Z,, 485, (1911) (36) s. H. Higgins, Chem. Soc. Proc., 29, 302, (1913) (37) E. K. Rideal and u. R. Evans, J. Soe. Chem. Ind. (Re• view), il, 64, {1921) (38) V. H. Remington and H. M. Trimble, J Phys. Chem.-, 33, 424, (1929) (39) E. K. Rideal, J. Chem. Tech., 1912, 141 (40) J. J. Weiss, z. Elektrochem., 37, 20 & 271, (1931) (41) J. D. Blakely, J. Soc. pyers and Col., !U!, 306, (1934) (42) s. H. Higgins, Chem. Soc. Proc., 28, 130, (1912) (43) s. H. Higgins, Chem. Soc. Proc., ~. 486, (1912) (44) A. Stock, Ber., ll, 837, (1920) (45) s. s. Perel 1 man and T. M. Lelyakina# Zavodskaya Lab., ll., 810, (1945)

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114 (._6) A. R. Choppin and L. c.Faulkenberry, J. Am. ~em •. Soe., fil!., 2203, (1937) {47) F. Dienert and.F. Wandenbuleke, Comptes rend., 1!, 29., ' ,, .., ; ,; C ' (1919) (48). s~ Gol~sc~~dt., Ber., g&. 753, (1919) (49). v. M. Clark, *The.Determination or Hydrogen Ions•, 3rd l ' I " ,l t (50) (51) (52) (53) Edition, Vill~ams.a~~ Wilkins, Baltimore, (~92~), pp •. 200-201 •standard Methods ror the.Examination or Water and ,\ ,, ,, ,, ,,. Sewage•, 9th Edition, American Public Heal~h Association, ~ew Y9rk, (~946), P• 98 Ibid, P• 152 ' ,, ' ' H. Bohm~ and.E. Schneider, Ber., 76 B, 483, (1943) D. s •. McKinney, Ind. E!Jg• Chem., Anal. Ed., !, . 192, .(1931) . '

PAGE 117

ACKNOWLEDGEMENTS This opportunity is taken to acknowledge vith grateful appreciation the many helpful suggestions and the encourage ment offered by Dr. A. P. Black, vho was Chairman of' the Supervisory Commi.ttee under vhieh this investigation was conducted. Acknowledgement is also made of' the valuable services rendered by Dr. A.H. Gropp, who showed a great deal of' interest in this work and on many occasions gave t'reely of' his time to f"urther its accomplishment, although he va.s not a member of' the Supervisory Committee. To him I express my l gratitude. It is further acknowledged that the Wallace and Tiernan Company offered a great deal of' financial assistance in the vay of' a f'ellovabip during the course of' this vork, and the opportunity is hereby taken to thank that organization for its help. 115

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BIOGRAPHICAL ITEMS . . . . JAMBS BROWN COODSON 1 JR. 1 was born July 10# 1915 1 in SandersYille,eeorgia. He accompanied his f'amily to Plorida in 1923 and attended, the public schools ot Gainesville. After &r&du.ating f'rom Gaineffille High School in 1933 he entered the Un1Yers1t7 of' Florida. where he received a Bachelor ot Science , < ~, degree 1n chemistry with honors 1n 1937. He round immediate employment as a chemist at the Georgetown,South Carolina, mill or the International Paper Company. In January of 1941 he was called to active duty 1n the Army or the Uni.ted States as a Second Lieutenant in the ChemicalYartare Service, and soon after-reporting t'or duty at Edgewood Arsenal. Maryland, he was assigned to a Chemical Laboratory Company. He commanded this organization With the rank or Major f'or more than three years,. the greater part of which were spent overseas in the Med1.terranean and Pacific Theaters. Be was discharged f'rom the Army early 1n 1946 vi.th the rank of Lieutenant Colonel and returned to the Inter• national Paper Company as Assistant Chief Chemist. He re-entered the University of Plorida 1n September of 1946 to pursue a course of studies leading to the Ph.D. degree in Sanitary Chemistry, with minors in Publie Health Engineering and Bacteriology. Mr. Goodson. is a member or Gamma Sigma Epsilon Chemical Fraternity and the American Chemical Society. 116

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This dissertation was prepared under the direction or the Chairman or the candidate's Supervisory Committee and has been approved by all members ot the Committee. It was submitted to the Graduate Council and vas approved as partial f"ulfilment or the requirements for the degree of Dooto~ of Philosophy. July 22, 1950 SUPERVISORY COMMITTEEz Cha1rman a21.dt4!= 117