Nov., 1960
1663
KINET~CS OF THE HYDROLYSIS OF CHLORINE
THE KINETICS OF THE HYDROLYSIS OF CHLORINE. I. REINVESTIGATION OF THE HYDROLYSIS I N PURE WATER BYASSALIFSHITZAND B. PERLMUTTER-HAYMAN Department of Physical Chemistry, Hebrew Universitg, Jerusalem, Israel Received April 25, 1960
In a reinvestigation of the kinetics of the hydrolysis of chlorine in pure water, using the thermal method, the experimental results of revious authors are confirmed. The two possible reaction mechanisms are discussed, and it is shown that OHcannot t a f e part in the rate-determining step.
Introduction In a recent investigation of the hydrolysis of bromine in phosphate buffer solution' we found that the observed rate was compatible with the assumption that the buffer anion takes part in the ratedetermining step. However the reaction was too fast for measurement by the flow method and therefore this assumption could not be quantitatively verified. The question of the mechanism according to which the homonuclear halogen molecule reacts to yield ionic products seemed however sufficiently interesting to warrant further research. Instead of trying to find the conditions under which the hydrolysis of bromine can be investigated, we thought it more promising to turn our attention to the hydrolysis of chlorine, firstly because in the case of chlorine the equilibrium
ever, to Shilov and Solodushenkov's second paper.2b Furthermore, the value of kz calculated by Morris is extremely high. In spite of this, his mechanism has been accepted, for instance, by SidgAs a first step in our program of investigating the kinetics of the hydrolysis of chlorine we decided to reinvestigate the reaction in pure water with the aim of clarifying its mechanism. Experimental
The rate of reaction was measured by the continuous flow method, utilizing the temperature effect. The apparatus used was essentially as described previously.1.6 An observation tube was employed which accommodated 12 thermistors. Temperature differences could be measured with a precision of j~0.001". The chlorine solution was prepared by bubbling gaseous chlorine (Matheson) under atmospheric pressure into one of the reaction-vessels while it was isolated from the rest of the apparatus by suitable stopcocks. The concentration of the chlorine water and of the reaction mixture was determined by measuring iodometrik, cally the concentration of the liquid as it ran out of the obXp Hz0 If HOX H + X(1) servation tube. k,' When the undiluted chlorine solution was passed through the observation tube, a slight temperature drop was noticed (where X represents the halogen) lies much more along the tube. This is probably due to the fact that we t o t h e right than in the case of bromine, and sec- employed a chlorine solution which was almost saturated ondly, because the rate of hydrolysis of chlorine in a t atmospheric pressure: owing to the pressure gradient pure water has been investigated2 and found to be along the observation tube the solution becomes superand endothermal de-solution of chlorine may take in a range which is accessible to measurement by saturated place. This difficulty does not arise in the case of the more the flow method. dilute and therefore more strongly hydrolyzed reaction Shilov and Solodushenkov2 assumed reaction 1 mixture. The changes of temperature due to the reaction, to represent the actual mechanism, although in their A T , a t each thermistor were therefore calculated by comthe temperature of the reaction mixture with that first paperza the rate constants corresponding to paring of the pure water. The temperature of the chlorine solution this scheme showed a downward trend as the re- relative to that of the pure water was determined from action proceeded. On reinvestigationlZbthis trend measurements at the first three thermistors only. There did not recur, and was ascribed to some systematic was usually a slight temperature difference between the two solutions and the values of A T were corrected by adding this error. difference, multiplied by the dilution ratio. On the other hand, Morris3 pointed out that the A reaction temperature of 9.5' was chosen on the basis decrease of the rate constant corresponding to of the following considerations: the temperature should be mechanism 1 can be explained if we assume the low in order to (a) keep the reaction rate within a range accessible to measurement in our apparatus (even in buffer rate-determining step to involve the hydroxyl ion, solution where the reaction may be expected to proceed whose concentration decreases as the reaction pro- faster') and (b) increase the solubility of chlorine. On the other hand, there exists a chlorine hydrate which dissociates ceeds. He therefore suggested the mechanism according to kz C124H2O(s) 8Hz0 Cldgas) Clz OHHOC1 C1(2)
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Recalculating the results in Shilov and Solodushenkov's first paper on this basis, Morris indeed found much better constancy for the rate constants. Morris' argument does not apply,4 how(1) A. Lifshitz and B. Perlmutter-Hayman, Bull. Research Council Israel, AS, 166 (1959). (2) (a) E. A. Shilov and S. M. Solodushenkov, Compt. rend. acad. sei. U.R.S.S., 1, 96 (1936): (b) J . Phys. Chem. (U.JS.JS.R.), 19, 404 (1946); Acta Physicochim. U.R.S.S., 20, 667 (1946). (3) J. C. Morris, J . A m . Chem. Soc., 68, 1692 (1946). (4) JI .inbar and H. Taiibe, zbtd., 80, 1073 (lQ58).
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As soon as the partial pressure of chlorine over the solution becomes equal to the dissociation pressure, any further addition of chlorine causes the hydrate to precipitate in the form of small crystals. The formation of this hydrate must obviously be avoided, i . e . , the temperature must be high enough for its dissociation pressure to lie above atmospheric pressure. The pressure us. temperature curve is very steep' ( 5 ) N. Sidgwick, "Chemical Elements and their Compounds," Vol. 11, Oxford University Press, Oxford, 1950, p. 1213. (C) E. Giladi, 4.Lifshitz and B. Perlmutter-Hayman, Bull. Research Council Israel, AS, 7 5 (1969). ( 7 ) "International Critical Tables," Vol. V I I , hIcGrax-Hill Book Co., Inc.. New York, N. Y . , 1980, p. 2.53.
ASSALIFSHITZ AND B. PERLMUTTER-HAYW
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danger of chlorine from one vessel entering into the other and dissolving. This indeed occurred, but to a very small extent, and was corrected for. Furthermore, the partial pressure of chlorine over the chlorine solution decreaaes during the experiment, and the solution may become more dilute. The chlorine concentration was determined before and immediately after the reaction had been carned out. A very small change was noticed, and the mean value was applied for calculating the kinetic results. The fortunate lack of equilibration between the chlorine solution and the gas-phase may be ascribed to the short duration of the experiment (-15 min.). (The position is different from that of previous work3 where the as bubbled through the solution and special precautions had to be taken.)*
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Fig. 2.-The dependence of Q ( x ) on time, for 5 experiments, presented in the order of decreasing stoichiometric chlorine concentration. (Straight lines correspond to mechanism 1.)
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Fig. 3.-The dependence of Dz(x) on time, for a typical experiment (run 111). (A straight line would correspond to mechanism 2.) and cuts the point of our atmospheric pressure (690 mm.) at 9.0’. A driving pressure of about 400 mm. Hg was used, and the flow velocity was between 5.35 and 6.52 cc. sec.-l. The pressure was a ainl.6 applied over both vessels from a nitrogen cylinder. %hue, while the mixture of the two reagents is being run t h u g h the observation tube, the two vessels are in contact uia the gas-phase, and there exists the
The rate equation corresponding to mechanism 1is ds/dt = ki(a
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where a is the total stoichiometric concentration of chlorine in the reaction mixture, and x t the molarity of the reaction products at time t. This yields on integrationZ
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3% tan-’ 2xt + x e d 3 x . 2 4K1 + d 3 X 2 4Kl const. = Dl(x) const. (11)
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where Kl is the equilibrium constant of reaction 1 and xe the value of x when equilibrium is reached. The value of xo, the amount hydrolyzed at zero time, is equal to the equilibrium value of x in the chlorine water before dilution, times the dilution factor. We usedg KI = 2.23 X which was corrected9 for the influence of ionic strength by dividing by the appropriate value of the square of the mean activity coefficientlo of HC1. (The values of zo and ze were calculated from Kl corresponding to the ionic strength before dilution, and at the end of the reaction, respectively. In expression I1 the mean of the values of K 1 corresponding to the ionic strength immediately after dilution, and at the end of the reaction was used.”) The values of a lay between (5.1 to 6.9) x lO-ZM, those of $0 between (1.43 t o 1.68) x 10-2 M and those of xe between (2.07 to 2.38) x 10-2 M.12 The values of zt were calculated from the expression (8) Nevertheless, it a.ould have been more desirable to apply a driving-force of chlorine over the chlorine solution and a driving-force of nitrogen over the bottle containing pure water. However, the slightest fluctuation in one of the pressures would then cause a considerable change in the ratio of the amounts of liquid flowing from the two bottles. This effect can be greatly decreased by inserting between the vessels and the mixing chamber capillary tubes whose hydrodynamic resistance is a t least equal to that of the mixing chamber and observation tube. This device, however, cannot be applied. Apart from the fact that it necessitates the use of inconveniently high pressures to achieve the desired flow velocity, it entails a large pressure drop between the reagent vessels and the mixing chamber which would cause part of the chlorine to de-solve. (9) R. E. Connick and Yuan-tsan Chis, J . Am. Chem. Soc., 81, 1280
(1959). (10) H. 9. Harned and B. B. Owen, “The Physical Chemistry of Electrolyte Solutions.” Reinhold Publ. Corp., New York, N. Y., third ad., 1958, p. 716. (11) If the change of K Iduring the reaction were taken into account the integration of equation I would become very difficult. However, K I changes during the reaction only by about 4%: furthermore, expression I1 is not very sensitive to small changes in X i . The error introduced by neglecting the change in KI is therefore small in comparison with our experimental inaccuracy.
KINETICS OF THE HYDROLYSIS OF CHLORINE
Nov., 1960
1665
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Zt = 20 (Z. - ~ o ) A r t / A T r . ( 1111 Morris,* it leads to a rate constant of 5 X l O I 4 where ATt is the temperature effect at a given mole-' 1. sec.-l, and zero activation energy (the thermistor, and AT is the total temperature effect. increase of the rate of hydrolysis with increasing The measured values of AT- lay between -(44 temperature is ascribed to the change in the disto 50) X 10-3 degrees, and were in good agreement sociation constant of water). However, the highest with those calculated from (ze - xo) andQ AH1 possible value for a second-order rate constant has been shown's to be much lower than this, whereas = 6.55 kcal. mole-'. Figure 1 shows a plot of AT vs. for a typical a value of zero for the apparent activation energy experiment (run 111), and Fig. 2 shows the de- seems incompatible with the mechanism of difpendence of Dl(r) on time for 5 experiments. fusion-controlled reactions. An even stronger arguStraight lines are seen to be obtained which show ment against mechanism 2 is obtained when we no trend as the reaction proceeds. The values of kl consider the question whether OH- - which a t a can be calculated from the slope of these lines. pH of about 2 is present in the solution at a concentration very much lower than chlorine-can We find be regenerated at a sufficient rate. According to kl = 5.60 sec.-l Morris' estimatela this may indeed be the case, with a standard deviation of 0.45. The com- but in the meantime, both theoretical consideraparatively high scatter is probably due to the tions's and direct measurement,14 have shown that small value of (2, - zo) and therefore of the total a t 25" the rate of formation of OH- in water is heat effect. A small error in AT will therefore only about 1.4 X 10-8 mole 1.-l sec.-l, i.e., much cause a considerable error in D(x). (This is especi- slower than the rate of hydrolysis of chlorine (about ally true for the last few points of each experiment, 0.1 mole 1.-1 see.-' a t our concentration). This where xt approaches x e . ) excludes the possibility of the hydroxyl ion playing I n addition we plotted the expression (see also any important role in the hydrolysis of chlorine Morris,s eq. 6) in pure water. (Experimental results so far obtained by us show that the situation may well be different in buffer solution, where OH- can be supplied faster than in pure water.lJ4) The reaction thus proceeds according to the reaction scheme represented by equation 1. This against time. If mechanism 2 were operative, does not answer the question of the actual rnechathis would yield straight lines. Concave curves nism, namely whether we have to consider the rewere obtained. An example (run 111) is shown in action as a monomolecular one in which a hyFig. 3. drated chlorine molecule dissociates with a rate Discussion constant of 5.6 sec.-l or as a bimolecular one inThe fact that the plots of Dl(z) us. time exhibit volving collision between chlorine and water, and no downward trend as the reaction proceeds, proceeding at a specific rate of lo-' mole-' 1. set.-'. whereas the plots of D&) exhibit an upward trend At the moment, there seems to be no possibility of shows mechanism 1 to be compatible with the ex- deciding this point on theoretical grounds, but an perimental results. Apart from this experimental investigation of the rate of hydrolysis in buffer confirmation of mechanism 1, mechanism 2 can be solutions may be expected to throw some light on ruled out on theoretical grounds, which have been the question. The authors wish to thank Prof. G. Stein for stressed in recent years.ls Firstly, according to making val liable comments. (12) Unfortunately, we cannot employ any drastic changes in a, and o)
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hence in m and ze. since we are limited by the decreasing heat effect at low concentration, and the solubility of chlorine at high conoentration.
(13) See e.0. hf. Eigen, Disc. Faraday Soc., 17, 194 (1904), where the relevant literature is quoted. (14) M. Eigen and L. de hIaeyer, 2. Eleklrochem., 69, 980 (1955).
