Peroxdisulfate-Iodide Reaction Kinetics in DMSO: Mechanism and Salt Effects, Study notes of Law

Download Peroxdisulfate-Iodide Reaction Kinetics in DMSO: Mechanism and Salt Effects and more Study notes Law in PDF only on Docsity! AN ABSTRACT OF THE THESIS OF BRUCE BURNS for the MASTER OF SCIENCE in INORGANIC CHEMISTRY Date thesis is presented August 13, 1965 Title THE KINETICS OF THE REACTION BETWEEN PEROXYDI- SULFATE AND IODIDE IONS :IN DIMETHYL SJJLFOXIDE Redacted for Privacy (Major Professor) 6 Abstract approved Persulfate oxidizes iodide ion according to the equation 52082 + 3I --- 2SO42 + I3 The distinctive feature of the persulfate- iodide reaction in dimethyl sulfoxide (DMSO) is that it has.a two term rate law, viz: Rate = dl3 /dt = ka S2O8 ì+ kb[I i i S2O82-I. A mechanism has been presented in which the iodide independent oxidation is brought about by a radical intermediate which is formed when sulfate radical ions react with DMSO. The exact nature of this radical intermediate is not certain. Evidence for the non -existence of a direct interaction between SO4 T and I is given by the fact that the first -order reaction is not observed in water and that it dis- appears in a solvent mixture of 70% DMSO /30% H2O. The values of kb exhibit a positive salt effect which seems to I/ be dependent on total cation concentration rather than ionic strength. The data were found to fit the equation kb =(N)o + 2.16 x 10 -3[K1 where(k )o 3. 00 x 10 -5M -1 sec .1. The k a exhibits a negative salt effect. Tetraethylammonium perchlorate was found to have no effect on the reaction while barium nitrate exhibited a much greater accel- erating effect on kb than did potassium ion. = ACKNOWLEDGEMENT I wish to acknowledge my indebtedness to Professor J. H. Krueger for his assistance and encouragement during the course of my research and for his valuable advice on the writing of this thesis. TABLE OF CONTENTS I. INTRODUCTION 1 Effect of Solvent DMSO on the Reaction 2 First -Order Decomposition 3 Reaction of Persulfate and Iodide 7 Mechanism 7 Salt Effects 8 II, EXPERIMENTAL 12 General . . 12 Apparatus 13 Reagents 14 Experimental Procedure 17 Runs in which Persulfate was Analyzed +4 17 Runs in which Absorbing Species (I 3 or Ce ) were Analyzed 18 Calibration of DU 24 Treatment of Data 24 III, RESULTS AND DISCUSSION 33 Mechanism 33 The Effect of Oxygen 38 Effect of Added Cations 43 Solvent Effects 54 Effect on kb 54 Solvent Mixtures 56 Catalytic Effect of Chloride 58 Effect of Allyl Acetate 60 The Reaction of Per sulfate and Cerous Ions . , 60 IV, CONCLUSIONS 62 BIBLIOGRAPHY 64 LIST OF FIGURES Figures Page la. Plot of Ro/ S2082 vs. [I ]for calculations of ka and kb at constant u = 0.0600 M 27 l b. Plot of Ro/ S2081 vs. [I for calculation of ka and kb at constant g= 0. 0500 M. . . 28 lc. Plot of Roo S2082] vs [I] for calculation of ka and kb at constant 4 = 0.0400 M. . . . 29 ld. Plot of Ro/ S2081 vs. [I-] for calculation of ka and kb at constant k(= 0.0300 M. 30 -]1 e. Plot of Ro/ Is vs. [I-] for calculation of 1 ka and kb at constant 4= 0. 0200 M . . 31 41 45 46 47 -DMSO mixtures (100% -O, 85.00% -0, 70.00% -, 50, 00% -®DMSO by weight). . . 57 2. Comparison of runs done with (0) and without (0) nitrogen sweep at kt = 0.0500 M. 3. Log kb vs [K+;1 4. kb in DMSO vs. total potassium ion 5. Log kb in DMSO vs. square root of ionic strength 6. Plots of Ro/ S2082 vs. [I ] for various H2O 1 I 1 l ] 2 increased ion pairing between the reactant anions and added cations over what might be expected in water. In addition, studying the reaction in mixtures of water and DMSO would provide a means of isolating specific solvent influences which are bound to be present in an aprotic solvent. From preliminary investigations it was apparent that the reaction was zero or near zero order in iodide ion although iodine was produced as a product of the reaction. This suggested a very important solvent effect. Later experiments, however, done in the absence of oxygen showed that the iodide concentration did affect the rate of the reaction but with an order less than one. This indicated that the persulfate- iodide reaction in DMSO had a two term rate law just as the persulfate- bromide reaction has in aqueous so- lution (9, p. 182). In a later section it will be shown that a mecha- nism can be written which leads to a rate law consistent with the data and which involves a thermal decomposition step for persulfate. Effect of Solvent DMSO on the Reaction The data obtained in this study for the persulfate- iodide reac- tion have brought several facts to light which are focused in three areas: 1) The rate law in DMSO involves two terms including an iodide independent term which is absent in aqueous solution; 2) The rate of the second -order reaction is slower in DMSO than in water; 3) Cation effects manifest themselves in a somewhat 3 different manner than in water. These are the three factors which differentiate these results from those obtained in aqueous solution. Thus, any conclusions to be drawn about solvent effects of this reaction must account for these differences. The next two sections will briefly summarize the findings of other workers on the first -order decomposition of aqueous persul- fate and the persulfate- iodide reaction. First -Order Decomposition Aqueous persulfate solutions are known to decompose slowly at somewhat elevated temperatures. The decomposition proceeds via two paths, one of is catalyzed, the rate law being Rate = -dS2082 /dt = kl +k2IH1IIS2O82] (23), Several steps have been proposed as the initial and rate determining step for the un- catalyzed decomposition, (12), viz! S2082 --a. 2SO4 S2082- 2SO4 S2082 SO4 + S042- (2) (2a) (2b) Reactions (2a) and (2b) provide paths for the exchange of sulfur between S2082 and S35 labeled SO42 The mechanism for (2a) involves the entirely reasonable equilibrium SO4 + S*04 - S042- S*04. . which . + , 4 No such exchange has been observed (7, 29), and, thus, these steps are considered unlikely. This initiating step (reaction 2) may also be brought about by light, gamma rays, or impurities or dust in the solution (12). In reactions with organic substrates and water reaction (2) is rate determining so that the observed rate is always first -order with respect to 52082 and zero -order with respect to the oxidizable substrate except at very low concentrations. In water the initiation step (2) is presumably followed by reaction with the solvent to pro- duce OH' radicals SO4 + H20 --a. OH- + HSO4 . (3) If an oxidizable substrate is present in the solution it can be attacked by the SO4 radicals or by the OH-,but the slow step re- mains the same. Yet it has been observed that the nature of the oxidizable substrate influences the observed rate constant for the reaction. Thus, the radicals produced when SO4, radicals attack the substrate probably influence the decomposition of the S2082 ion. Presumably, the S041- radical itself exhibits this same effect, yet this premise has not been verified (12). House (12) has written a general mechanism for persulfate decomposition in the presence of an oxidizable substrate which fits the organic substrates which have been reacted with persulfate; Reaction of Persulfate and Iodide Mechanism This reaction falls into a different class from the first -order decomposition type because it is also first -order in the reducing agent. In aqueous solution it is observed uncomplicated by any first -order decomposition because the decomposition is so slow. Three mechanisms have been proposed to account for its first -order dependence on iodide and persulfate: Reference 52082 + I -a IS2083- slow (28) (11) IS2083 + I -0. I2 + 2S0 2- fast (12) IS2083 -a I' + 82083 slow (31) (12a) 3 +I- -0,.. 2SO4- + I' fast (12b) I' + I' --f I2 fast (13) S2082 + I IS2O83- IS2083- -a I+ + 2SO4 2- + I + I --4. I2 . (12) Modern authors seem to favor reactions (11a), (lib), and (13a) as the correct mechanism principally because of the existence of the 7 (lla) (11b) (13a) 8 compound bis- (pyridyl)- iodine (1) persulfate. This compound is formed when persulfate oxidizes potassium iodide in pyridine. There is certainly a question as the exact nature of the complex IS2083 . The above mechanisms all indicate that it is a reactive intermediate. However, since it has not been isolated, it could as well be the activated complex. The two possibilities are not kinetically distinguishable, and the point is unlikely to be re- solved unless the IS2O83 species is actually isolated in a compound. Nothing further need be said about this mechanism since its import- ance in this laboratory centers around the fact that it has as a rate determining step a reaction of two anions S2082 and I . Salt Effects The Br$nsted theory of reaction rates which utilizes the Debye- Húckel limiting law gives the relationship between rate constant and ionic strength for an ionic reaction as 2Z Z, (( log k = log k + ^- log k + o 2ZAZB, o lßa where ko is the rate constant at /1= O, ZA and ZB are the charges of the reacting ions, a is the average distance of closest approach of two ions in solution, an4BandXare constants characteristic of a given solvent and temperature. The persulfate- iodide rate con- stant has been found to conform with equation (14) at low ionic 14) iss yµ 9 strengths. King and Jacobs (20) found that the rate constants agree with those predicted by the approximate form of equation (14) up to Vh,L= 0..16 providing the ratio of divalent to univalent ions in solu- tion was kept low. Log k 0 then was found to be -1. 075. However, in two sets of experiments where large concentrations of divalent ions were present log k 0 was -1.06. While these values are in near agreement there may be some significance in their differences. Negative deviations from the limiting expression were observed with MgSO4present ata= 0.06. Since it inherently contains all the assumptions of the Debye- Huckel limiting law equation (14)fand especially its approximate formlis valid only at very low ionic strengths and in the presence of only uni- univalent electrolytes. Therefore, it is scarcely surprising that equation (14) has been found invalid for a wide variety of ions at moderate ionic strengths. Using data of their own plus those of other workers Olson and S imonson (25) derived an equation which fits data for several reactions between ions of like charge. The equation is 1 " K(x) k = ka 1 +K(x) + k 1 +K(x) (15) where K and k are empirical constants, (x) is the total concen- tration of some ion, and ka is the specific rate constant of the reac- tion at some low ionic strength. If k a = k o (ie:k a= k o at x = o) the term in the brackets, may be considered a substitute for Brp&nsteda s a 12 II. EXPERIMENTAL General Since most of the experiments in this study were performed using pure DMSO as solvent it was considered imperative that water be excluded as thoroughly as possible. Since most reagents used were nonhygroscopic it was generally considered sufficient to dry them overnight in an oven at 110oC and store them in weighing bottles over CaSO4 or CaC12 as drying agents. This was the standard pro- cedure with all solids unless otherwise noted. Solid materials were weighed out on a Mettler balance of 200 gram capacity. In experiments where samples were analyzed by titration the burets used were Pyrex 10 ml class A burets. All pipets used to mix solutions or remove samples were calibrated by filling with freshly distilled thermostated DMSO, emptying into tared glass containers and weighing. Densities of pure solvent and solvent mix- tures were taken fromCowie and Toporowski (5). All pipets and other glassware involved directly in making up stock solutions or solutions for kinetic runs were cleaned with di- chromic acid cleaning solution and rinsed many times with distilled water. Most glassware was oven dried at 1100C. 13 Apparatus Two constant temperature baths were used during the course of experiments. One consisted of a ceramic crock of approximately three gallon capacity fitted with a "lightnin" model L mixer, a 250 watt blade type heater, and a mercury thermoregulator switch of standard design set at 19. 67 + 0.02°C. The temperature was check- ed with a 500 mm 50o thermometer. The second bath consisted of a styrofoam ice chest approxi- mately 10" x 12" x 18" fitted with a pump which circulated water through the jacket around the sample holder in the spectrophotometer, a 250 watt blade heater and thermoregulator set at 19.80 + 0. 03°C. Both baths were cooled by three turn coils of copper tubing at the bottom through which flowed ordinary tap water. The coolant water for the second bath was pre -run through about 10 feet of copper tubing coiled inside a small crock filled with ice water. This procedure was necessary in the summer since the tap water was 22oC, but was not used in the winter. Most of the experiments were carried out in the Beckman model DU Quartz spectrophotometer. The DU was thermostated as men- tioned above. The cells used were standard Beckman Quartz cells. The DU was powered by a storage battery which was charged contin- uously during the time the instrument was on at ca 2. 5 amp but only 14 a trickle (a few tenths of an amp) when the instrument was off. The instrument was operated at 6.4 volts. Reagents Dimethyl sulfoxide was obtained from Crown Zellerbach in polyethylene lined five gallon drums. Straight from the drum it was clear and had only a faint odor. It was guaranteed to contain less than one percent water, but for analysis of iodine solutions small concentrations of dimethyl sulfide or higher sulfides present the greatest difficulty. The solvent was distilled from an all pyrex glass system which included a 15 inch column packed with glass helices at a pressure of 10 -20mm. The distillate had a boiling point which was quite constant during a distillation but varied according to the pressure obtained from the water aspirator (70- 85 °C). Start- ing with about one liter of unpurified DMSO a middle fraction of 500 ml was collected after two distillations. Although the sulfide content was not checked there seemed to be no anomalous results such as rapid reduction of persulfate to cause concern. The water content in one sample was determined by Karl Fischer titration to be 0.03 %. Potassium dichromate was obtained reagent grade and recry- stallized from water before being dried and stored. It was used as a primary standard. 17 solid was always handled in a dry bag filled with dry nitrogen until a sample was weighed into a small, dry stoppered volumetric flask anddissolved in DMSO. Two solid samples were analyzed by oxidiz- ing the Ce(III) and DMSO by boiling with K2S2O8 in excess and back titrating with standard ferrous solution using ferrous 1, 10 phenanth- roline as indicator. Samples were reoxidized and retitrated until two successive determinations gave molecular weight values within 0. 5 %. The average value of the molecular weight was 546 +2 grams/ mole corresponding to a formula of Ce(NO3)3. 2. 82 DMSO. Tetraethylammonium perchlorate was prepared by precipitation from aqueous solution according to the method of Kolthoff (22). Eastman tetraethy lammonium chloride was dried in vacuo over P2O5 and used without further purification. Experimental Procedure Runs in which Persulfate was Analyzed Runs five and ten were done in_solutions containing no iodide, and, hence, were not done in the DU. The object of these runs was to determine the rate of reduction of per sulfate by the solvent. The procedure in both experiments was to dissolve a known amount of K2S2O8 in DMSO, place it in the thermostated bath,and remove samples from time to time for analysis. Unfortunately, the 18 persulfate is slow to dissolve in DMSO and the use of the magnetic stirrer to get it into solution in a reasonable amount of time was necessary. This caused a great deal of initial decomposition because the stirrer had quite a tendency to heat up the solution. In ten, for example, the K2S2O8 concentration dropped by about 10% from the time the solution was mixed to the time the first sample was with- drawn. Therefore, zero time was taken as the time of withdrawl of the first sample. Time was kept with an electric timer and read to nearest second. The time of withdrawal of a sample was taken as the time when half the pipet had emptied into the quench solution. The quench was composed of 10. 00 ml of standard ferrous solution, 50 ml 1N HL SO4 solution and 5 ml concentrated (85 %) phosphoric acid. The samples were back titrated with standard dichromate and the per - sulfate concentration calculated by difference. No attempt was made to exclude oxygen from the solution since oxygen has no affect on the first-order decomposition. Runs in which Absorbing Species (I3 or Ce +4) were Analyzed The solutions in the remainder of the experiments contained either iodide or cerous ion both of which yield products which can be analyzed by their absorption of light in the near U. V. -far visible region. In these experiments the products were analyzed and it was most convenient to allow the reaction to procede right in the quartz 19 cell in the DU. The appearance of iodine (as triiodide)was followed by observing the change in optical density of the solutions at 370 mkt ; the cerium (IV) at 340 m (,( . The solutions were made up by weighing the desired amount of solid into a 25 or 50 ml ground glass stoppered flask, weighing, adding approximately the desired volume of solvent, and reweighing. In some earlier runs the K2S2O8 solution was weighed into the reac- tion vessels, but in later runs the KI solution was weighed. All the other solutions were pipetted from Pyrex 5 ml graduated pipets or a Pyrex 1 ml graduated pipet. When these pipets were allowed to drain in a dropwise manner the volume delivered corresponded to the volume read to within one or two parts per thousand so the dif- ference was ignored. It is clear that the molar concentrations of each of the solutions which was pipetted was in error by some small amount because they were made up by weight and delivered by volume. Inherent in the method is the assumption that the molarity and molality of these solutions are equal. Since most stock solutions of K2S2O8 and KI were .02 to . 03 M the error in this assumption is probably fairly small (ca. 0.5 %). But the more concentrated solutions of KNO3 (0. 1 to 0. 2 M) were considerably more in error (up to 4 %). Gen- erally, this would affect the ionic strength so that the(,( values were least accurate of all the concentrations. Unfortunately, these errors 22 this line, then, is equal to kb and the intercept at I I = O is equal to k . a The fact that such plots did give a straight line with non - zero intercept verifies the above rate law since the reaction has been shown to be first -order in persulfate. Two experiments were performed in the presence of barium ion from Ba(NO3)2 to demonstrate the expected greater rate accel- eration of Ba ++ over K+ and one experiment was done in the pre- sence of tetraethylammonium perchlorate which was selected as an example of a cation which would exhibit little or no ion pairing with 52082 , and, therefore, have less influence on the rate. Three experiments were done with allyl acetate added to the reaction mixtures. Allyl acetate was expected to act as a radical trap for the SO4 T radical ions and thus "isolate" the second -order reaction by preventing iodine from being formed from SO4 T ions. Allyl acetate concentrations ran from 0.01 to 0.12 M. Two experiments were done using Ce(III) as the reducing agent in an attempt to "isolate" the first -order decomposition. The Ce(III) was added as the DMSO -solvated nitrate in concentrations of 4. 5 x -3M " -3 and 1.0 x 10 10 M at varied persulfate concentrations. Several experiments were done in various solvent mixtures of DMSO and water from 85.00 %to 50.00% DMSO by weight in order to establish the effect of the DMSO on the reaction rate as a specific or general one. These were done using exactly the same procedure as J 23 the ones in pure DMSO but at only one total potassium ion concen- tration (0.0450 M). A final experiment was done with tetraethylammonium chlor- ide in the reaction mixture to discover a possible catalytic effect on the decomposition by choride ion. In addition the results of one experiment (no. 13) are included which were not used in any calculations. This experiment is pre- sented to verify the fact that the reaction is first -order in per sulf- ate. The rate constants thus obtained were not used in any calcula- tions because the experiment was done in the presence of atmos- pheric oxygen. Because this reaction is so slow it must be studied by an ini- tial rate method. Since this is true, each kinetic run yields only one value of Ro and 5208 2 at one initial [5208 2 and I 1 Thus the plots in Figures la through le represent many runs. This is somewhat of a disadvantage because less information is obtained from each experiment than may be obtained from a faster reaction. In addition, small differences in the experimental technique may influence the self- consistency of the data. Throughout this research every effort was made to adhere strictly to a standard procedure in setting up and making the runs. I 24 Calibration of DU The extinction coefficient of I3 in DMSO was determined at 370 mg. Standard 13 solutions were made up by weighing re- agent grade resublimed iodine into weighed portion of DMSO. Fur- ther dilutions were done by weight until final concentrations of were (0.8 to 3.0) x10 -5M. The final dilutions were made in 10 ml volumetric flasks and diluted with spectroscopic grade DMSO which was also 0. 05 M in KI. Thus, the high ratio of I to I2 assured complete conversion to I3 -. The extinction coefficient of I3 in DMSO was determined to be (2,16 + 0, 01 )x104 Beer's law is obeyed. Treatment of Data It is quite common in studying reactions involving free radi- cals to obtain scattered values for the rate constants. This is because free radical reactions are often catalyzed by minor impuri- ties in the solution, dust, imperfections in the glassware, and other gross mechanical factors which are hard to control. Clearly this is true of the first -order persulfate decomposition. Referring to Table (I) and Figures (la) through (le) it is apparant that the data are scattered. In most cases this scatter is probably attributable to either the factors listed above affecting the rate of the first -order reaction or by the presence of oxygen in the solution which reduces I3 0 4 5 [I] x103M Figure la, Plot of Rot 2082 vs. [I jfor calculations of kaand kbat constant ,U.= 0.0600 M. l Figure lb. Plot of Ro 6 7 8 9 IO II 12 [IJxIO3M S208 2 vs. {1]for calculation of ka and k at constant (,(= 0.0500M. o L6 14 L2 O 4 5 [I1xIo M IO Figure lc, Plot of Ro/ Is 2082 1 us.II l for calculation of ka and kb at constant L 0.0400 M. /J = 32 the rate of the second -order reaction. Thus, it is not surprising that about half the points in each plot fall off the best straight line. The lines were drawn by the method of least squares, those points being eliminated from the calculations which were more than (1 %) away from the line drawn by eye through the points. In figure 1 d there are actually two straight lines of different slope and intercept which could be drawn. The indicated line was selected as the "best" because more points fell on it than the line of greater slope. The points denoted by ft5 were not used in the least squares calculation. The same procedure was used for the studies at other solvent compo- sitions but the requirements were relaxed somewhat for the data in 85% DMSO because there were too many "borderline cases". In the 70% and 50% solvent mixtures so few points were available that they all had to be used. The same is true of the runs done with Ba(NO3) 2 added. 33 III. RESULTS AND DISCUSSION Mechanism There appear to be three reactions occurring simultaneously in DMSO solutions of KI and K252O8 The first is the oxidation of DMSO by persulfate. The second and third produce triiodide and re- sult in the two term rate law Ro= dl2 = ka [S208 2 + k I [2082_] dt (21) The first reaction, oxidation of solvent by S2O8 2- , does not yield a detectable product when solutions are analyzed for I3 at 370 mg , but it must occur. Since the pure DMSO used in the various experi- ments contained 0.03% water, only one water molecule was present for every 1000 DMSO molecules. Since water is tightly bound to DMSO in solution, as evidenced by their obviously high heat of mix- ing as well as the non -ideal behavior of DMSO -H2O mixtures (5), it is unlikely that any water molecules are available to react with the shortlived SO4' radical. Therefore, reaction (4a) is not possible. So a mechanism must be written which is different from the sulfoxide- persulfate mechanism in water. A comparison of experiments five and ten in Table II gives evidence that the persulfate decomposition in the solvent is at least roughly first -order in persulfate. The dif- ference in [ S2O821 values (which are really first -order rate 1bi a[ J J L Table II. Reduction of per sulfate by DMSO and KI Run LKZSZO8]x103M [KI]x 103M total i((x 103M Rox 107sec 1 Ro/ 52082 x lObsec 1 5 12.28 24.56 36.84 0.680 5.53 10 93.2 186.4 279.6 5.50 5.96 11 77.4 36.8 181.6 269.0 2.80 3.63 103M L o 0 8 -- 37 to one of two things: (1) the iodide ions and DMSO molecules are competing for the sulfate radical ions or (2) the iodide ions are not reacting directly with sulfate radical ions at all but are competing with the sulfate radical ions for the DMSO derived radical inter- mediate designated above as 03S0- S -O(CH3)2. Of course, there is always the possibility that both of these competitions are going on. Unfortunately, this confusing set of possibilities cannot be resolved by the data presented here. There is one fact which suggests that the second alternative is the correct one, namely, the first -order persulfate /zero -order iodide reaction is not observed in aqueous soltuion. This suggests that DMSO is indeed catalyzing the reaction by, for example, the following mechanism: k SO4- + : (CH3)2S0 --2.. O3SO S(0)(CH3)2 (21) O3SO S(0)(CH3)2 + I k- SO4 + (CH3)2S0 + I. (22) I + I. -+ I2 (23) Any solvent molecules which react with sulfate radical ions but do not react with iodide are not significant since they produce no de- tectable products. 2- The third reaction is one in which I and S208 react directly. The mechanism for this reaction in aqueous solution has already been discussed. There is no reason to believe that the mechanism 38 will be any different in DMSO. The data in Table III show that the overall reaction is first - order in persulfate, since the basic rate law is Ro= ka + kb I [s2c8 2- s (21) dt = [S2082-] Ro/ S208 2 = ka + kb I (26) which is a constant when iodide is constant as it is in experiment 13. The numbers themselves are, however, qualitative because the experiment was done in the presence of oxygen. The Effect of Oxygen Table IV contains the results of two experiments done in the presence of atmospheric oxygen. Figure 2 shows Ro+2082 values from Table IV plotted versus I I . The values for the rate constants are ka= 0.48 x 10 -6 and kb= 0.70 x 10 -4 These values may be compared with the data from experiments 33, 35, 37 and 38 which are also plotted in Figure 2 and which yield values of ks= 0.510 x 10 -6 and kb = 1.30 x 10 -4 The effect on kb is mark- ed while the effect on ka is less important. In addition it should be noted that values of Ro/ S2O8 2 I tended to vary a great deal at constant I 1 when oxygen was present, possibly due to slightly varying amounts of 02 actually in an individual solution. The I I I I . . -1 -I a Table III. Verification of first -order in persulfate. Run x10 [KIjx 103M [Klt otalx 103M ,c{x 103M Ro/ {s2082} x 106 sec -1 13 A 1. 58 20. 2 42. 3 43. 9 2. 50 B 3.16 20.2 40.7 43.9 2.66 C C 4.75 20.2 39.1 43.9 2. 59 D 6.33 20.2 37.6 43.9 2. 29 E 7.91 20.2 36.0 43.9 2.41 42 results reported in Table IV are the most self consistent obtained without sweeping the solvent with nitrogen, The small difference in the above noted ka values is probably not significant. It should be remembered that all of the errors in the runs are thrown onto these values, and they are probably only good to within + 10% in general. In particular, Figure 2 shows that the scatter of points about the line is such that the ka intercept could be higher than the intercept of the N2 sweep line. The kb values are, on the other hand, strikingly different. Examination of the mechanisms listed above (equations 11-13) sug- gests no clear cut mechanism by which oxygen could interfere with the production of iodine. It is well known that oxygen has the ability to act as a radical scavanger because it has two unpaired electrons. For the same reason it may enter into a reaction as the propagator of a radical chain. Thus the rates of some catalyzed per sulfate oxidations are increased by oxygen (33). The fact that oxygen does have such an important effect on this reaction is strong evidence that it is a radical chain mechanism. Radical chain mechanisms may be written for the second -order reaction, but all of them envolve either a sulfate radical ion or something very much like it. If oxygen was a scavenger for SO4r a significant effect on the first -order reaction would be observed. Since this is not the case, the explanation for this oxygen inhibition a 43 must remain unaccounted for. Effect of Added Cations In Table I are listed the results of experiments done at various ionic strength values with the iodide varied. Figure 1 shows the Ro/ [s2082 values from this table plotted versus I I for calcula- tion of kb and ka values. These values are listed in Table V. Figure 3 shows a plot of log kb versus[ Kitotal which is corn- pared with a plot of Perlmutter- Hayman and Stein's log kb values calculated from the A, B, and D values they obtained for K+ (see equation 18). Clearly if the data from DMSO is to agree with equa- tion (18) the values of the constants must be somewhat altered from their values in aqueous solution. This, in itself, is no draw back to the use of equation (18), but reference to Figure 4 shows that a plot of kb versus [Kir is linear as predicted by equation (15). Thus equation (15) seems to be a better equation to describe the re- lationship between kb and cation concentration for the range investi- gated than equation (18). The product of constants kak a K from equation (18) is calculated to be 2.16 x 10 -3 from the slope of the line in Figure 4. Figure 5 shows a plot of log kb versus which is also a straight line as predicted by the Br$nsted -Debye equation. The slope of this line is 3.36 which determines the value of a in L l 1 1 / 44 Table V. Values of ka and kb at various ionic strengths. ( x 103M [Kitotal x 104M kbx 104M lsec 1 kax107 sec 60. 0 55.0 1.55 4.47 50. 0 45.0 1.30 5.10 40. 0 35.0 1.10 7.52 30. 0 25.0 0. 862 5.43 20. 0 16.3 0. 662 8. 00 b a 42 CA J -4 -44 -45 -4. .06 .10 .14 .18 NIT M12 Figure 5. Log kb in DMSO vs. square root of ionic strength. -3.8 -39 -4.0 -41 48 equation (14) to be 0. 840 in contrast to a value of 0. 509 in water. The values of aCemployed in these experiments are somewhat high- er than those usually expected to obey equation (14). Amis and Potts (1) did obtain agreement with the approximate form of equation (25) at values of .212 and above. The value of of is related to dielectric constant and tempera- ture by the following equation: Q(= K (DT) -3/2 Since the value ofp( in water at 25°C is equal to 0. 509 the value of a in DMSO at 20°C can be readily calculated, (( 80.4 x 298) 3/2 = 2.02 .509 48.9 x 293 ( =1.03 The values of D in water and DMSO being 80.4 (1) and 48.9 (30) respectively. The true value of of is higher than that calculated by equation (14). This apparent disagreement can readily be explained as will be shown below. The effect of added barium ion on the reaction rate is even more striking than that of potassium ion as would be expected. Only one experiment was done with barium ion added (Table VI). The value of kb at a total L = 0.0199 M and Ba ++ = 4.16 x 10 -3 M was 2. 00 x 10 -4; in contrast to a value of 0.738 x 10 -4 from Figure 4 at the same total with no added Ba ++ . Using - [K+] (27) ( Table VI a. Added Ba(NO ) Run 57 I J K ( ; 3 3 jK2S208,x 10 M [Klj x10 M iK l 4.96 0.913 4.96 3.170 4.96 5.271 3 Ixl0 M 19.9 19.9 19.8 ++ Ba 4.16 4.16 4.16 3 10 M R 2- S2O8 6 -1 x 106 sec 0. 667 1.08 1.56 Table VI b. Added Et4NC1O4 (65) and Et4NC1 (66). Run K2S2O8x 103M Klfx 103M `K+totalx 103M [Et4Nlxlo3M Ro lObsec -1 J ( 1 _ 65 0 5.00 6.16 P 5.00 6.16 Q 5.00 6.16 66 R 5.03 6.13 S 5.03 6.13 T 5.03 6.13 34.8 5.79 34.8 11.58 34.8 17.37 35.0 6.88 35,0 13.07 35.0 19.27 , 1.30 1.30 1.30 0.524 0.491 0.431 ( - J S2O82 tl l 52 reaction at all but only the specific effect of something like ion pairing. In other words, the linearity of log kb versus Vµ noted above is coincidental. The above results do not weaken the argument about ion pair- ing. The Et4N + ion has a very small ionic potential. It was chosen for study because it is an ion which would be expected to ion pair very little with S 0 2 8 2.- . According to Table VI it exhibits no ion pairing at all. The ka values listed in Table V also show some dependence on ionic strength. In this case, however, the ka value decreases with increasing ionic strength. It happens that three of the values (at /J = 0.06, 0. 05, and 0. 02) fall on a straight line when log k a is plotted versus /with a slope of -2. 38. The interpretation of why this effect is observed is a great deal more difficult than in the second -order reaction because the rate determining step in this reaction involves only one ion, 5208 2 If, for example, the effect is due to ion pairing it might be true that the KS2O8 ion, whose charge is somewhat localized on the oxygens farthest from the potassium, would be less likely to break apart than the 5208 2 ion whose charges are more or less localized at opposite ends of the ion. a a , Thus, the reactions are - O O O O 53 I I I I unpaired O-S-O-O-S-O -0-S-0-+O-S-0- I 1 I 1 O O O O O O O O I I I I paired KO-S-O-O-S-O- KO S-O + O-S-O. I I I I O O O O The electrostatic repulsion between the ends of the unpaired ion would serve to enhance the transition state and products. But with one end of the ion paired this repulsion does not exist. But it is difficult to estimate the magnitude that this effect may be expected to have. In addition, there may be other effects which require a 2- more intimate knowledge of the transition state and the S2O8 ion than is available at present. Nevertheless, the effect does exist to a significant extent despite the large errors which may exist in the listed ka values, and it seems possible that further study of the salt effect on ka might lead to important discoveries about the first - order reaction. When Kolthoff and Miller (23) studied the thermal decomposi- tion of aqueous persulfate solutions, they found that, while the un- catalyzed rate constant was unaffected by added salts, the acid cata- lyzed rate constant exhibited a negative salt effect. This fact casts some doubt on the ion pairing argument advanced above as a 54 possible explanation for the salt effect on ka in DMSO. Since, if both the k reaction in water and DMSO have the same rate deter - mining step, ion pairing should influence both reactions about equal- ly, Kolthoff and Miller advanced no explanation as to why their reaction had a salt effect. The acid catalyzed decomposition mechanism advanced as: H+ + S2O82 HS2O8 HS2O8 ---o SO4 + HSO4 SO4 - S03 + 1/2 02 (stoich. ) strong SO4 + H2O H2SO5 (detectable) acid (30) (31) (32) (33) seems to bear little relationship to the decomposition mechanism proposed above for DMSO solution. Solvent Effects Effect on kb Strictly speaking, two of the previous sections(Mechanism and Effect of Added Cations) have dealt with solvent effects. Already then, it is apparent that there are differences between the persul- fate- iodide reaction in water and DMSO. Another obvious difference a --s. 100% 50°/ 0 3 6 9 [I] x 103 70% 12 15 18 Figure 6. Plots of Ro/ [s2082J vs. [I, for various H2O -DMSO mixtures (100% -G, 85.00% -0, 70.00% -I, 50.00% -®DMSO by weight). 3 as more water is added to the solvent. No ready explanation presents itself to explain this phenomenon. The surprising k a 58 and kb values in the 50.00% DMSO by weight mixture suggest that something was radically wrong with the experiment. The solvent effect on k a , however, is more reasonable. Re- ference to Figure 6 will show that ka seems to decrease rapidly from 100% DMSO to 7.0.00% DMSO where it practically disappears. This indicates that the k a reaction does not occur unless the con- centration of water is very low. This supports the contention made above that the species which reacts will iodide ion in this mechanism comes from DMSO and that I does not react with SO4_ directly. Catalytic Effect of Chloride It was thought that chloride ion might exhibit a catalytic effect on the second -order reaction which does not exist in aqueous solution. To discover this effect, one experiment was done with varied amounts of tetraethylammonium chloride in the solution. The results may be found in Table VII. All these values for Ro/ 52082 are lower than those taken fromFigure lc which has the same [K +] total value. This is due to two factors: (1) the data in Table VII were obtained at 380 mg instead of 370 mil which was the usual wavelength; (2) the chloride ions compete with iodide ions a Table II. Reduction of persulfate by DMSO and KI Run wt. %[K2s208]x DMSO 103M [KI]x 103M [K+lttlx x10 3 M Ro/ S2O82 x 106M 1 sec 1 59 0 85,00 4.97 1.028 44. 8 0.320 P 85.00 4.97 4.487 44. 8 0.413 Q 85.00 4.97 8.210 45. 0 0. 951 61 U 85.00 4.98 3.473 45. 0 0. 362 A 85.00 4.98 10.453 45. 0 0. 682 B 85.00 4.98 14.275 45. 1 1.02 62 C 85.00 4.95 8.006 45. 1 0. 543 D 85.00 4.95 9.196 45. 1 0. 645 E 85.00 4.95 12.63 44. 7 0. 835 68 I 70.00 5.04 3.920 45. 0 0. 204 J 70.00 5.04 11.027 45. 2 0.672 K 70.00 5.04 15.367 44. 8 0. 974 64 L 50.00 4.99 4.379 44. 8 0.816 M 50.00 4.99 12.871 45. 1 0.946 N 50.00 4.99 20.846 45. 0 1. 10 Table VIII. The reduction of Ce (III) by per sulfate. Run [Kzszo82ixlo3M {Ce (III) x103M ,( x 103M Ro/ S2O82 1 2M-1 sec -lx E 34 A 1.287 4.50 49.9 0.94 B 3.753 4.50 50.3 0.980 C 6.172 4.50 50.0 0.979 56 F 3.520 1.09 60.1 0.026 G 9.049 1.09 60.2 0.0222 H 11.19 1.09 59.8 0.0278 *Data uncorrected for extinction coefficient of Ce (IV) in DMSO at 340 mkt o t 62 IV. CONCLUSION The most remarkable facet of this reaction is the two term nature of the rate law in sharp contrast to its strictly second -order nature in water. This feature has already been discussed in some detail and it is to be regretted that so little data is available with which to determine more accurately the nature of the k a reaction. It seems that valuable answers could be obtained by isolating the products of the persulfate decomposition in pure DMSO. Also more data is needed on the relationship between solvent composition, ionic strength and k a . Further study of the first -order reaction could yield valuable information on the nature of all first -order persulfate oxidations. The tentative conclusion that the second -order reaction is in- fluenced by specific cations only and not by all cations generally, even at these moderate ionic strengths,is one which is in need of further verification. It is not surprising to find deviations from a frankly limiting equation like the -Debye equation, but the fact that tetraethylammonium ion had no effect at all on the rate is startling indeed. It is unfortunate that more experiments were not done at various concentrations of multivalent ions like Ba ++ or La + ++ to firmly nail down the salt effect as being one of ionic strength or cation concentration only. Another feature is the slowness of the second -order reaction 63 in DMSO as compared with water. It would be interesting to see if other reactions between ions of like charges behave in a similar fash- ion in DMSO. This would serve to test the conclusion that highly charged transition states are destabilized by lack of hydrogen bonding, The question as to the extent to which ion pairing influences this reaction is one which cannot, as yet, be answered. It clearly must play some role in DMSO even with its moderate dielectric con- stant, especially in the case of di -and trivalent cations. 'In general, DMSO isa satisfactory solvent in which to study ionic reactions. DMSO has the advantage of having the highest dielec- tric constant of any aprotic solvent. Most commonly employed water soluble salts are soluble in DMSO to some extent. Exceptions to this rule are chlorides, sulfates, and other salts containing anions of high ionic potential. Thus, kinetic studies of ionic reactions can be made which yield interesting solvent effects associated with lack of hydro- gen bonding, etc. 66 21. King, C. V. and Eric Jette. The oxidation of iodide ion by persulfate ion. II. The effect of removing the products of the reaction on the reaction velocity. Journal of the American Chemical Society 51: 1048 -1057. 1929. 22. Kolthoff, I. M. and J. F. Coetzee. Polarography in aceto- nitrile. I. Metal ions which have comparable polarbgraphic properties in acetonitrile and water. Journal of the American Chemical Society 79: 870-874. 1957. 23. Kolthoff, I. M. and I. K. Miller. The chemistry of persulfate. I. The kinetics and mechanism of the decomposition of the persulfate ion in aqueous medium. Journal of the American Chemical Society 73: 3055 -3059. 1951. 24. Marshall, Hugh. C ontributions from the chemical laboratory of the University of Edinburg. no. V. The persulphates. Journal of the Chemical Society 59: 771-786. 1891. 25. Olson, A. R. and T. R. Simonson. Rates of ionic reactions in aqueous solutions. Journal of Chemical Physics 17: 1167 -1173. 1949. 26. Pederson, Kai Julius. The effect of metal ions on the rate of decomposition of nitroacetic acid. Acta Chemica Scandinavica 3: 676 -696. 1949. 27. Perlmutter- Hayman, B. and G. Stein. Specific ionic effects on reaction rates. The reaction between persulfate and iodide ions in the presence of high concentrations of added salts. Journal of Chemical Physics 40: 848 -852. 1964. 28, Price, Thomas Slater. Die Reaktion zwischen Kaliumpersulfat und Jodkalium, und Katalyse bei der selben. Zeitschrift für physikalische Chemie 27: 474 -512. 1898. 29. Riesebos, P. C. and A. H. W. Aten, Jr. Absence of rapid exchange of sulfur atoms between sulfate and persulfate ions. Journal of the American Chemical Society 74: 2440. 1952. 30. Schlafer, H. L. and W. Schaffernicht. Dimethylsulfoxyd als Losungsmittel für anorganische Verbindungen. Angewandte Chemie 72: 618 -626. 1960. 67 31. Soper, Frederick George and Emyln Williams. The effect of the solvent on reaction velocity. III. The interaction of per - sulfate ions and iodide ions. Proceedings of the Royal Society of London 140A: 59 -70. 1933. 32. Winstein, S. and Stanley G. Smith. Sulfoxides as nucleophiles. Tetrahedron 3: 317-321. 1958, 33. Woods, R., I. M. Kolthoff and E. J. Meehan. As (IV) as an intermediate in the Fe (III) and Cu (II) catalyzed As (III) - persulfate reaction. Inorganic Chemistry 4: 697 -704. 1965.