+ IO3-, at a rate proportional to [H+]2[Br Thirdly, the iodine bromide reacts by the overall

Transcription

1 Reaction between iodine and sodium bromate D. E. C. KING AND M. W. LISTER Depnrtmet~t of Chetilistry, Utliversity of Toronto, Torot~to, Otimrio Received July 18, 1967 The mechanism of the reaction: I2 + 2BrO3- --, Br in aqueous solution (ph 1.5 to 2.5) at 25 "C has been investigated. It is found to go in four stages. Firstly, there is an induction period during which the acidified bromate produces a catalyst (probably HOBr) for the next stage. However, iodide ions, if present, destroy this catalyst. Secondly, iodine reacts by the overall equation, 12 + Br IBr + IO3-, at a rate proportional to [H+]2[Br Thirdly, the iodine bromide reacts by the overall equation, 31Br + 2Br H Br ,- + 6H+, at a rate proportional to I[Br:l2PrO3-]2. Finally bromide and bromate ions react to give brom~ne. Mechan~sms are suggested for the various stages, and some of the rate constants are evaluated. Canadian Journal of Chemistry, 46, 279 (1968) Iodine is known (1) to displace bromine from and Lomb Spectronic 505. The other analytical methods aqueous bromate solutions, although, of course, consisted of extracting the reacting mixture with carbon tetrachloride, and estimating the halogen extracted the is the other way around in either spectrophotometrically, or by addition of aqueous halide solutions. The total reaction, which is potassium iodide and titration with sodium thiosulfate. Iodine-thiosulfate titrations were also used to standardize I2 + 2Br03--+ BrZ , the stock solutions of iodine and sodium bromate. The substances used were all of analytical reagent grade, can be expected to go in a number of steps, usually dried overnight at 115 OC, and made up by weight since it is unlikely that so many oxygen atoms in standardized volumetric glassware. would be transferred in a single stage. It seemed The stock iodine solution might be mentioned in a therefore to be of interest to investigate the little more detail. British Drug Houses Analar reagent mechanism of this reaction, especially as it iodine was shaken with doubly distilled water 10 to 15 times, and these solutions were discarded. In this way might be expected to yield more information on it was hoped to remove impurities such as traces of oxy-halogen reactions in general. As will be iodides or of lower halogens. The final solution was made seen in what follows, the reaction did indeed up by shaking a few small crystals of iodine, which had prove to be complicated, and although the been treated in this way, with fresh water. The saturated mechanism has been elucidated to a considerable solution so produced, which is about 10-3 ilil, was extent, only rather speculative proposals can be put forward for certain parts of the reaction. Experimental The experiments in general consisted of mixing aqueous solutions of iodine and sodium bromate, with added phosphoric acid and sodium dihydrogen phosphate as a DH buffer. The DH of the reaction mixture was in the ;ange of 1.5 to i.5; and was measured on a Beckman Research ph meter using a glass electrode, Beckman 40498, in the usual way. The ph reading was calibrated by means of a potassium tetroxalate buffer (ph 1.68) but various dilute HC1 solutions of ph about 1.0 to 1.5 were checked against this buffer. It was assumed that the HC1 was completely ionized, and consequently these measurements gave a correction which enabled ph readings to be converted to hydrogen ion concentrations. Sodium nitrate was also added to the reaction mixtures to control the ionic strength, which was always close to The reacting solutions were kept in a thermostat at 25.0 =!= 0.1 "C. The reaction was chiefly followed by measurements of optical density on a Beckman DU. spectrophotometer. In some cases spectra over the range 200 to 500 ml were measured on a Bausch poured off, standardized, and used as the stock solution. In a typical run various amounts of the reacting solutions, the buffer, and sodium nitrate, initially at 25 "C, were rapidly mixed, and the reaction followed as described above. In almost all mixtures the bromate was in considerable excess, typical initial concentrations being: [I?] = 6 X 10-4 M, pro3-] = 2 X 10-2 M. Results The reaction was found to occur in four well-defined stages. (i) An induction period, during which very little, if any, iodine reacted. The length of this period depended on the concentrations, but was typically 20 to 30 min. (ii) Iodine disappearance. This lasted usually 1 to 2 min, and during this time iodine was replaced by iodine bromide. (iii) Iodine bromide disappearance, This lasted usually 2 to 5 min. At the end the solution was colorless, and contained sodium bromide. (iv) Bromine appearance. This lasted several

2 CANADIAN JOURNAL OF CHEMISTRY. VOL. 46, 1968 TABLE I Values of the slope, s = d log D-]/dtl, for various concentrations Series [H+l (M) pro3-] (M) Slope Kl (min-1 M-3) hours, and eventually gave a bromine concentra- slowly producing a product which catalyzes the tion equal to the initial iodine concentration. reaction of iodine and bromate ion. This These stages will now be considered in detail. catalyst, however, reacts with traces of iodide ions which are present, and it is only when 1. Induction Period virtually all the iodide has been used up that the Initially the color of the solution, as estimated iodine reaction can begin. The evidence for by eye, stays constant for some time, and then these statements will now be given. suddenly fades. It was found that the length of The absorption spectrum of a reaction mixture this induction period could be estimated was found to be as shown in Fig. 1. Initially the visually to within a few seconds (roughly 0.05 absorbance followed curve (a) with maxima at min). Spectrophotometric measurements at ,355, and 290 mp. The peak at 460 mp is due mp, which detect iodine with very little interto iodine, and the two other peaks are attributed ference from iodine bromide, also gave constant to triiodide ions, which are reported (2) to have absorbances during this period. Experiments maxima with extinction coefficients of at were also made in which the reacting solutions 355 mp, and at 290 mp. These last two were shaken with carbon tetrachloride, and the maxima slowly diminished giving curve (b) iodine so extracted was estimated, either from after some minutes. If a trace of potassium iodide the absorbance of the CC14 layer at 510 mp or was then added, it restored the absorption curve by shaking the separated CC14 layer with approximately to curve (a). On further standing aqueous potassium iodide and titrating with the triiodide peaks disappeared again at the sodium thiosulfate. These experiments showed same rate as before. no loss of iodine within experimental error The induction period was, as might be expected, during the induction period. lengthened by added potassium iodide, and the The existence of an induction period was extent of this lengthening was investigated. attributed to the following mechanism. The If TI is the induction time, a plot of log (added acidified bromate ions are believed to decompose iodide concentration) against TI was linear for larger amounts of added iodide. This linearity FIG. 1. Absorbance during the induction period of a solution initially containing 2.0 x 10-2 M NaBr03 and 6.0 x 10-4 M I2 at ph could be extended to all the measurements if - it - was assumed that the concentration of iodide ion present as an impurity was 1.5 x 10-5 M (in 6 x 10-4 M iodine), and log (total iodide) was plotted against TI. The concentrations of added potassium iodide varied up to 2 x 10-4 M. The slope, d log [I-]/dTl, was measured at various ph and bromate ion concentrations with the results in Table I. These results can be explained by supposing that the slow stage is the reaction I- $ Br H+ -t products, and that the induction period ends when [I-] has fallen to some critical value. The above equation requires that -d ln [I-] - kl[~+]'[13roy-] cl t

3 K~NG AND LISTER: REACTION BETWEEN IOD~NE AND SODIUM BROMATE 281 where kl is the rate constant. To relate kl to the slope, s, in Table I, a correction must be made for the fact that much of the iodide ion is combined as 13-. The equilibrium constant for is reported (3) to be 725 M-1, and as [Iz] was 6.0 x 10-4 M in tliese runs, this means that 69.7 % of the total iodide was free I-. Hence the relation of s and kl should be Values of kl obtained from this equation are given in the last column of Table I. The mean value of kl is 1.87 x 104 min-1 M-3. Earlier workers on this reaction find the same rate law as given above, with values of kl of 2.9 x 104 min-1 M-3 (4), or 3.5 x 104 min-1 M-3 (5). These values were for somewhat different ionic strength, but if we consider the rather indirect way in which kl is obtained in the present work, the agreement is sufficiently good that there can be little doubt that removal of iodide is responsible for the induction period. Extrapolation of the results to zero induction time suggests that this would correspond to a concentration of iodide ions of about 4 x 10-6 M, which is presumably the concentration at which iodide ions cease to compete effectively with 6 x 10-4 M iodine for the decomposition product of the acidified sodium bromate. Various other substances were added in order to find their effect on the induction period. Sodium bromide had no effect, though it did on the next stage of the reaction. Sodium iodate reduced, but did not totally remove, the induction period. This is somewhat surprising in view of tlie explanation offered above, but perhaps the reason is that though iodate and iodide ions react at ph 2, iodate ions do not help to produce the catalyst required in the next stage of the reaction. It might be added that kinetic data on the reaction of iodate and iodide ions, measured by Abel and Stadler (15), show that the explanation of our observations is not that tlie reaction is quite slow at the low iodide concentrations in the present experiments. Sodiunl broinite, prepared as in ref. 16, at a coilcentration in the reaction mixture of 1.5 x 10-4 M, reduced the induction period to about 2 mill. Hypobromous acid solution (prepared fro111 aqueous broiniile and silver nitrate) at about 10-4 M removed the induction period completely. In this preparation, bromine was in a slight excess, so bromine as well as hypobromous acid was added: however, it was found that small amounts of bromine alone did not greatly reduce the induction period. The bromine would presumably mostly react to give iodine bromide. Hypoiodous acid (prepared from aqueous iodine and silver nitrate and very rapidly filtered into tlie reaction mixture) at about 10-4 M reduced the induction period to less than a minute. Solutions of hypoiodous acid, though unstable, can exist long enough to have the kinetics of their decomposition studied (6); and it can be seen that hypoiodous acid behaves differently from its decomposition products, iodine and iodic acid. The hypobromous and hypoiodous acid solutions are saturated with silver bromide and iodide respectively; however, these silver salts, by themselves, did not reduce tlie induction time. It was suggested earlier that acidified sodium bromate slowly produces a catalyst for its subsequent reaction with iodine. The effect on the induction period of various substances described above suggests that this catalyst is either bromous or hypobromous acid, illore probably the latter, since broinous acid produces liypobromous acid during the course of its decomposition (7). Possibly both are active, though there is some balance in favor of hypobromous acid, since this removed the induction period most strikingly. Another set of experiments was made in which the sodium bromate was acidified and allowed to stand for some time before the iodine solution was added. It was found that tlie triiodide ion absorption peaks, mentioiled above, were immediately reduced after the addition, showing that some material was accumulating in the sodium bromate, which oxidized iodide ions rapidly. If To is this "pre-acidification" period, and Ti the induction period, it was found that a plot of log TI against To was approxin~ately linear, provided To was not too long. Thus we can write log TI -- n - DTo, where n and b are constants. The values of the slope, b, each of which was obtained from a number of experiments, are given in Table 11. It can be seen that b depends on ph and [Br03-1. Both tlie sodium bromate and iodine solutions

4 CANADIAN JOURNAL OF CHEMISTRY. VOL. 46, 1968 TABLE I1 Variation of the slope, b = -d log Tl/dTo with ph and [Br03-] Run PH [BrO3-1 (M) b* b/[h+l2[bro~]2 *~ime in minutes. were brought to the ph values given in Table 11, before mixing, so the ph was constant throughout the experiment. The results in Table I1 show that at any broinate concentration, b is roughly proportional to [H+]2. The bromate concentrations in Table I1 are for the "pre-acidification" period, and these were halved on adding the iodine. Between 4 and 2 x 10-2 M bromate, b is proportional to [Br03-12, but b is lower than would be expected on this basis for 1 x 10-2 M sodium bromate. The last column in Table I1 gives values of b/[h+]2[br03-12, obtained with the help of the correction to the ph (0.115 to get -log [H+]) mentioned above. These observatioils are tentatively explained by the following proposals. (i) The catalyst, X, is produced at a constant rate in any one run, though this rate may depend on ph and [Br03-1, and it also decomposes so that its concentration is always low. (ii) When the iodine is added, the catalyst rapidly reacts with some of the iodide ions present, and is thereby destroyed. (iii) The remaining iodide ions then react with the acidified bromate during the induction period, by the usual iodide-bromate reaction. When the iodide concentration is low enough, the catalyst builds up again. (iv) The dependence of b above on the concentrations probably indicates that the production of catalyst involves two hydrogen and two bromate ions, perhaps by the reaction (as a first slow stage) in particular, the mode of decon~position of the catalyst is unknown. Either a first or second order decomposition gives a curve of T1 against To of the right general shape, and the extent of the reduction of TI for not too large values of To depends on the rate of production of the catalyst, not on its decomposition. 2. Disappearance of IorEine At the end of the induction period the rate of disappearance of iodine increased rapidly. As the iodine absorption peak vanished it was replaced by new pealts at 390 mp and 255 nlp. The absorptioil at 410 mp remained constant for a considerable part of the time of disappearance of the iodine, so that it behaved rather like an isosbestic point. This presumably indicates that there is a constant ratio between the amount of iodine disappearing and the amount of the new colored compound appearing. The rate of disappearance of iodine was followed at 500 mp, where there was very little absorption by the new colored intermediate. The rate of decrease of absorbance accelerated rapidly at first, but then stayed constant until the iodine had almost all disappeared. This constant rate was measured for different conditions with the results in Table 111. The slope in Column 2 of this table is -da/dt, where A is the absorbance at 500 mp. Column 5 shows that the slope divided by [H+]2[BrO3-12 is reasonably constant, with an average value of 1.20 x 107 min-1 M-4 (average deviation x 107). This means that It is not possible to put these proposals into -- mathematical form with any confidence, since, -d dt ["I - ~ ~[H+I~[B~o<]~

5 KING AND LISTER: REACTION BETWEEN IODINE AND SODIUM BROMATE TABLE I11 Slope Run Slope (rnin-1) [H+l (MI IBrO3-I (MI [H+]2[Br03-12 and as the light path was 1 cm, and the extinction coefficient of iodine at 500 mp is 466, it follows that k2 is 2.6( 0.2) x 104 min-1 M-3. The colored intermediate was found to be iodine bromide, the evidence for this being as follows. Firstly, the solution was shaken with carbon tetrachloride, approximately at the instant when the iodine had disappeared, and gave a red solution (in CC14) with an absorption peak at 496 mp, which agrees with a peak previously reported for iodine bromide (8). The concentration of the iodine bromide was determined by shaking the carbon tetrachloride solution with aqueous potassium iodide, and titrating with sodium thiosulfate. From this the extinction coefficient at the maximum was calculated to be 394; the literature value is 390. Aqueous iodine bromide was prepared by mixing equimolar amounts of iodine and bromine solutions. As the equilibrium constant for is 8.3 x 104, from free energy data (9), the formation of IBr is nearly complete. The extinction coefficients of aqueous iodine bromide were obtained from this solution, after corrections were made, firstly for hydrolysis by the reaction and secondly for the reaction IBr + Br- = IBr2-. Mean 1.20 Data from ref. 9 make the equilibrium constant of the first reaction M9, and of the second 370 M-1 (10). The ph of the solutions were measured after adjustment to somewhere in the range of ph 1.0 to 2.0 by addition of a little nitric acid, and the concentrations of all the species present were calculated. The extinction coefficients of IBr2- were obtained by adding sodium bromide at a concentration of 0.05 M, when most of the iodine bromide is converted to IBr2-. The results are given in Table IV, together with those for iodine. The curves for iodine and iodine bromide cross at 412 mp, close to the point where the absorbance stayed constant during most of the disappearance of iodine. This indicates that at this stage of the reaction one iodine molecule is being replaced by one iodine bromide. This was confirmed by the observation that the amount of iodine bromide extracted into carbon tetrachloride was sometimes as high as 83 % of the maximum amount calculated from the equation It was also found that addition of sodium bromide (about M) to the reaction inixtures when the iodine had just disappeared gave the characteristic absorption peak of IBr2- at 375 mp. The reported (1 1) extinction coefficient at 375 mp is 590, in agreement with our results, and this was used to calculate the concentration

6 CANADIAN JOURNAL OF CHEMISTRY. VOL. 46, 1968 TABLE IV Extinction coefficients (E) of various species Wavelength (mp) E of IBr E of IBr E of I TABLE V d[ibr]-l/dt Run (M-1 min-1) [H+l (M) [Br03-1 (MI s/[bro 3-12 of IBr in the reaction mixtures. These were about 90% of the maximum amount calculated from the equation above. Presumably a little iodine bromide has reacted before all the iodine has disappeared. The mechanism of this stage of the reaction will be considered briefly later. 3. Disappearance of Iodine Bronzide The disappearance of iodine bromide was followed by its absorption at 390 mf. It was found that a plot of (absorbance)-1 against time was linear. As tlle extinction coefficient of IBr is 346 at this wavelength, tlle slopes of these lines could be converted to values of d[ibr]-lldt, and these are collected in Table V for experiments under different conditions. The bromate ion concentrations given in Table V allow for the amount of bromate consumed by the reaction I? + BrO3- -+ IBr + IO3-, the initial iodine concentrations being about 5 >: 10-4 M. The results in Table V show that d[ibr]-lldt is proportional to [Br03-12 and the last column gives the slope d[ibr]-lldt divided by [Br There is also probably some dependellce on ph, since tlle last column increases as [H+] decreases, but it does not seem to depend on a simple power of [H+]. However, the ph reading fell a little during this stage of the reaction, as would be expected since the reaction produces hydrogen ions. The phosphate buffer limited the change of ph to about Consequently we may write with kt about 1.5 x 108 min-1 M-3, but somewhat dependent on ph. As before, the mechanism will be discussed later. 4. Appearance of Bromine The equation for the removal of iodine bromide was given above, but so far the only evidence for it has been the observation that the ph fell somewhat during that stage of the reaction. Further evidence comes firstly from the observation that when the iodine bromide had disappeared tlle solution was colorless to the eye, and secondly from the data on the last stage of the reaction which will now be described. The appearance of bromine was followed by its absorption at 390 mp, and it was found (as was already known (1)) that the final concentration of bromine agreed within experimental error with the overall equation The measured absorbailces were found to be consistent with the following explanation. The total reaction up to tlle start of this final stage is It was assumed that the concentration of bromide ions at the beginning of this final step in the

7 KING AND LISTER: REACTlON BETWEEN lodlne AND SODIUM BROMATE TABLE VI Values of -d log [Br-]/dt for various conditions Run -d log [Br-]/dt (min-1) [H+l (M) [Br03-I (M) k x x x x OO OO OO 1.52 Run with added sodium bromide Part Added [Br-] (M) -d log [Br-]/dt (min-1) a 0.0 x x 10-3 b c rl reaction was determined by this equation, and that their subsequent concentrations could be calculated from the amount of bromine present in accordance with the equation Hence [Br-] could be calculated at various times, and it was found that a plot of log [Br-] against time was linear. Experiments were carried out at various ph and [Br03-1, and in addition some runs were made in which known amounts of sodium bromide were added to the reacting. u mixture. In the latter case it was found that log [Br-] against time was still linear, with the expected slope, thus confirming that the calculated amount of bromide, formed by reaction of the iodine, was indeed correct. Table VI gives values of the slope -d log [Br-]/dt for various conditions, and it can be seen that the slope divided by [H+]2[Br03-] is approximately constant, as shown in Column 5. This column gives the calculated rate constant of the reaction, defined by -d ln [Br-I - k3[~+]2[~r03-]. clt In the run with added sodium bromide, the calculated [Br-] at the beginning of part (a) was 1.16 x 10-3 M, the bromate concentration was 2.0 x 10-2 M, and hydrogen ion concentration about 3.6 x 10-3 M. It can be seen that the calculated d log [Br-]/dt stayed constant. The results in Table VI make k3 = 1.33 x 103 min-1 M-3 wit11 an average deviation of 0.08 >: 103, at 25 "C and an ionic strength of Literature values are 0.58 x 103 (5) at an ionic strength of 0.5, 1.5 x 103 at 0.013, 0.95 x 103 (12) at 0.21, and 0.65 x 103 min-1 im-3 (13) at These are a little lower than our results, but Young and Bray (12) extrapolate k3 to 2.7 x 103 min-1 M-3 at zero ionic strength. It may be noted that earlier values of the rate constant are generally reported in terins of d[br03-]/dt, which is one fifth of d[br-]/dt. The agreement between our rate constant and these earlier values is sufficiently good that they support the view that this is the same reaction. Discussion of Results This section will be an attempt to explain the results given above by mechanisms for the various stages of the reaction, and will be necessarily rather speculative in places. The various stages will be considered separately. I. Ind~~ction Period The experiments seen1 to establish the fact that the induction is due to reaction of iodide ions, partly with bromate ions, and partly with a species which develops in the acidified sodium bromate and which catalyzes the next stage of the reaction. The kinetics of the reaction of bromate ions with either bromide or iodide suggest that the slow stage of the reaction is (where X- is bromide or iodide), with a reactive intermediate XBr02 of the sort proposed by Taube and Dodgen (14). This type of intermediate will be suggested for later stages of the present reaction.

8 \ ~ ~, 286 CANADIAN JOURNAL OF CHEMISTRY. VOL. 46, 1968 The initial stage of the production of the catalyst mentioned above is probably 2H+ + 2BrO HBr , in agreement with the kinetic data, but as bromous acid would decompose further (7), 2HBr02 -+ HOBr + Br03- + H+, it is possible that either bromous or hypobromous acid could be the catalyst. 2. Reinoval of Iodine The overall reaction in this period is Iz + Br IBr , which proceeds over most of its range at a constant rate (in buffered solutions with a large excess of bromate ion) proportional to [H+]2 [Br It is difficult to suggest a mechanism which does not require the rate to tail off at the end more than is observed. However, the following proposal Rate cotistatits 2H+ + 2Br03- --, 2HBr k, 2HBr02 4 HOBr t- H+ + BrO3- k~, HOBr + I2 4 HOI + IBr k, HOI + Br03--+ HI02 + BrOz- HIO. + Br BrO2- + H+ leads to the equations Fast Fast where x = [I2], z = [HOBr], and R is the rate, assumed constant, of the first reaction. These equations were integrated numerically, and give curves of different shapes depending on the ratio of kc to k,. If kc is much the same as kb, a curve of [I2] against time with marked tailing off is obtained; if kc is very much greater than k, a curve is obtained in which the rate of disappearance of iodine increases continuously. If kc is about look,, a curve with a long nearly linear portion and very little tailing off is obtained, approximately what is observed. It is not possible from the present data to evaluate k,, and all that can be said about kc, therefore, is that it must be at least lookb. 3. Removal of Iodine Bromide Since this stage was apparently fourth order in the reagents, it seems reasonable to suppose that a rapid pre-equilibrium occurs, IBr + Br03- = BrIOBrOz-, to give a species rather like a polyhalide ion, and that two of these react slowly to give further intermediates, which in turn react rapidly to give eventually the final products. The slow stage is perhaps a transfer of oxygen, which is in effect a disproportionation. 2BrIOBr BrI-OBrOz- + BrIOBrO- I 0 then BrI-0Br02- I 0 + IBr -+ BrIOz + BrIOBrO- BrIOBrO- -+ Br- + BrI02 BrI02 + H20 -+ Br H+. This mechanism is, of course, only a suggestion to show that the final products can be reached by a fairly simple scheme. The species BrI02 is formally similar to other XY02 intermediates proposed for halide-halate reactions (14). 4. Al,yenrance of BI-omine This is believed to be the usual bromidebromate reaction with the slow stage Br- + Br H+ -+ Br.BrO2 + Hz0 followed by Br- 4- BrBrOz -+ Br2 + Br02- and other rapid steps which give bromine as the final product. 1. R. LYDEN. Finska Kemistsamfundets Medd. 37, 20 11,- 928). - --,. 2. A. D. AUTREY and R. E. CONNICK. J. Am. Chem. SOC. 73, 1842 (1953). 3. G. DANIELE. Gazz. Chim. Ital. 90, 1068 (1960). 4. R. H. CLARKE. J. Phvs. Chem (1906). 5. A. SKRABAL and H. SCHREINER. ~onatsh. 65, 213 (1935). 6. M. L. JOSIEN. Bull. Soc. Chim. France, 301, 814 (1948). 7. P. ENGEL, A. OPLATKA, and B. PERLMUTTER- HAYMAN. J. Am. Chem. Soc (1954): B. PERLM~ER-HAYMAN and G. STEIN. J. P&S. chem. 63,734 (1959). A. E. GILLAM and R. E. MORTON. Proc. Roy. Soc. London, Ser. A, 124, 604 (1929). W. M. LATIMER. Oxidation potentials. 2nd ed. Prentice-Hall. Inc.. New York D. 59. J. H. FAULL. ' J. dm. Chem. Soc. 56,-522 (1934). A. E. GILLAM. Trans. Faraday Soc. 29,1132 (1933). H. A. YOUNG and W. C. BRAY. J. Am. Chem. Soc. 54, 4284 (1932). J. SIGALLA. J. Chim. Phys. 55, 758 (1958). H. TAUBE and H. DODGEN. J. Am. Chem. Soc. 71, 15. E. ABEL and F. STADLER. 2. Physik. Chem. Leipzig, A122, 49 (1926). 16. Societe d'etudes Chimiques pour L'Industrie et L'Agriculture. Brit. Patent No. 843,558 (Aug. 4,1960).