ISOTOPIC EXCHANGE STUDIESBY ELECTRODEPOSITION decomposition of excited CHaT*molecules formed by an initial substitution of energetic tritium atoms into CHI. I n the propyl chloride systems, the single-step process is responsible for 97% attributable to the
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secondary decomposition of CaHaTC1*by elimination of HC1. Consequently, we conclude that none of the recoil tritium experiments to date offers positive support for the postulate of the existence of measurable yields of the single-step double-substitution reactions.
Isotopic Exchange between p-Nitrobenzyl Chloride and Chloride Ions as Studied by Electrodeposition1 by P. Beronius and H. Johansson Division of Physical Chemistry, University of Umeb, Umeb, Sweden Accepted and Transmitted by The Faraday Society
(April ??8,1967)
An electrodeposition method for following the exchange of chlorine between organic and inorganic chlorides is described. It has been used to study the isotopic exchange between p-nitrobenzyl chloride and lithium radiochloride in anhydrous acetone over a temperature ~ A comparison is made range from 15 to 35”. The exchange is found to be of S N type. with data for the same exchange reaction in aqueous acetone.
Introduction The application of electrochemical separation to the determination of precise rate data for bromine2 and iodinea-7 exchange reactions of the type RX X*- F?1 RX* X- has been described previously. Kinetic data for the isotopic exchange between p-nitrobenzyl chloride (p-N02BzC1) and chloride ions in acetone presented in this paper reveal that electrochemical separation is also a suitable means for following chlorine exchanges.
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Principle of the Method The electrochemical method used to follow the exchange reaction p-N02BzC1 36Cl- p-N02Bz36C1 C1- is analogous to that previously employed to study bromine2 and iodine3-’ exchanges. Small fractions of the inorganic chloride are anodically deposited on silver electrodes after different times of exchange and the activities of the silver radiochloride deposits determined. Since the activity of the deposit per unit amount of electricity is proportional to the specific activity of the inorganic chloride a t the reaction time in question, it’ is possible to calculate the rate of exchange. Experimental Section Equipment. A schematic picture of the reaction vessel (15-ml volume), which also served as an electro-
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lytic cell, is shown in Figure 1in ref 4. It consisted of a cylindrical aluminium vessel tightly closed with a lid of Teflon. The inner diameter and height of the vessel were 28 and 26 mm, respectively. The inside of the vessel was coated with Kel-F, a plastic fluorocarbon material, to prevent corrosion. It was provided with a water jacket connected to a constant-temperature bath controlled to within f0.02” using a certified thermometer. The easily exchangeable anode (see pp 9-12, ref 8) consisted of a 0.25 mm thick circular silver plate of 4mm diameter, vertically mounted at a distance of 9 mm from the centrum line of the cell. Only the front side of the anode (0.126 cm2 area), which had been carefully cleaned in a hot solution of a liquid detergent, was exposed to the solution. Its edges and back and the copper wire to which it was soldered were covered with (1) This Article is part X in the series “Electrochemical Methods in Kinetic Studies of Isotopic Exchange Reactions;” part IX, ref 2. (2) P. Beronius, Radbchim. Acta, 8, 57 (1967). (3) P.Beronius and 0. Lamm, Trans. Faraday Soc., 56, 1793 (1960). (4) P. Beronius, Acta Chem. Scand., 15, 1151 (1961). (5) P. Beronius, 2.Physik. Chem. (Frankfurt), 40, 33 (1964). (6) P.Beronius, ibid., 42, 45 (1964). (7) P. Beronius, Svensk Kem. Tidskr., 76, 646 (1964). (8) P. Beronius, Trans. Roy. Inst. Techml. Stockholm, No. 213 (1963). Volume 79, Number 2 February 1968
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P. BERONIUSAND H. JOHANSSON
.'"r I '
d ti t o 20 0 0
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t
8
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12
' ' "1
16
20
c.d. mA/cm2
Figure 1. Current efficiency as a function of current density and concentration for anodic deposition of chloride on silver electrodes from solutions of LiCl in acetone agitated at 400 rpm: electrolytic cell, 15-ml; [LiCI], 0.00030 M (open circles) and 0.0027 M (filled circles); temperature, 25.0'; anode area, 0.126 cm2.
wax (mp 69-73") to prevent electrodeposition. The wax was applied using a gently heated glass rod. After electrolysis, the wax was removed by immersing the electrode in petroleum ether (bp 60-71") for 0.25 hr. A platinum wire served as cathode. After each experiment this was cleaned by heating in a flame. The reaction vessel was completely filled with solution agitated at 400 rpm by means of a Teflon-coated magnetic bar (7 X 26 mm). This was rotated by an external magnet mounted on the axis of a synchronous motor. The electrolyses were carried out with constant current using a current stabilizer. The electrolyses times (18 sec in expt 1 and 27-36 sec in expt 2-9) were determined using a stop clock graduated to 0.01 sec and provided with a solenoid break. By this means, it was possible to arrange the circuits in such a way that the clock could be started and stopped at the moment the electrolysis current was switched on or off (see Figure 1,ref 8). Reagents. Acetone (pro analysi grade) was dried over anhydrous calcium sulfate and fractionally distilled in an atmosphere of dry nitrogen. This treatment reduces the water content to 0.001% or lessSg The purity of the solvent was checked by density measurements using a Lipkin pycnometer.'O The value 0.78433 g ml-l, obtained at 25.0", is to be compared with the literature value" 0.7845 g ml-I. The p-nitrobenzyl chloride used (Hopkin and Williams Ltd.) was recrystallized six times from absolute ethanol and dried in vacuo at room temperature. Anhydrous lithium radiochloride (Li3'jC1) with a specific activity of 9.6 mCi g-l, supplied by the Radiochemical Centre, Amersham, England, was dissolved in dry acetone. The exact amount of carrier chloride in this sample was determined by isotopic dilution analysis.12 Inactive lithium chloride (pro analysi grade) was dried in vacuo at 110" for 2 hr and stored in a desiccator. Solutions of lithium radiochloride in anhydrous acetone were prepared at 25". They were protected from the moisture of the air. The Journal of Physical Chemiatry
Kinetic Procedure. A weighed amount of p-NOT BzCl was dissolved in 15 ml of lithium radiochloride solution (0.1-0.5 pCi s6Cl),thermostatically maintained at the actual reaction temperature. The solution was transferred without delay to the reaction vessel and electrolyses were performed a t different reaction times. After each deposition, the anode was immediately removed from the electrolytic cell to minimize redissolution of the anodic deposit (cf. ref 13). The fraction of inorganic chloride deposited on each anode was 1% in expt 1 (Figure 2) and 2% in expt 2-9 (Tables I, 11). To minimize the disturbing effect of the analyses on the exchange reaction in the latter experiments, two analyses were made immediately after the reaction mixture had been transferred to the cell and another two analyses after approximately two half-times of exchange. The /3 activities of the silver radiochloride deposits were measured in a proportional counter. The standard deviation of the counting rate was reduced to k0.501, or less. In addition to the usual corrections, the activity values were corrected for the disturbing effect of the analyses on the course of exchange.14 The corrections did not amount to more than a few tenths of 1%. Calculation of the Rate of Exchange. The rate of exchange, R, was evaluated by a least-squares treatment of activity-time data using the equation16 a + c Rt log (1 - F ) = -ac 2.303 ~
where a = [p-N02BzC1]; c = [LiCl]; F = fractional exchange = (1 (c/a))(l - ( A / A o ) ) ;A. andA = electrode activity per unit amount of electricity at zero time of exchange and at time t respectively. Except in expt 1, no direct determination of A0 was made. This quantity could, however, be evaluated by successive approximations as described in ref 5.
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Results and Discussion Electrolysis Conditions. The electrolysis current was chosen to be within the optimum current density range for the electrode reaction Ag C1- 4 AgCl e-; cf. ref 8. To establish this optimum range, the dependence of the current efficiency on the current density was studied. Two examples of current efficiency-
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(9) A. Weissberger, Ed., "Technique of Organic Chemistry, Vol. VII, Organic Solvents," Interscience Publishers, Inc., New York, N. Y., 1955,p 382. (10) A. Weissberger, Ed., "Technique of Organic Chemistry, Vol. I, Physical Methods of Organic Chemistry," Interscience Publishers, Inc., New York, N. Y., 1959,Part I, Chapter IV. (11) M.B. Reynolds and C. A. Kraus, J . Am. Chem. Soc., 70, 1709 (1948). (12) P. Beronius, J . Electroanal. Chem., 9, 473 (1965). (13) P. Beronius, Trans. Roy. Inst. Technol. Stockholm, No. 234 (1964). (14) P. Beronius, 2. Physik. Chem. (Frankfurt), in press. (16) H. A. C. MoKay, J . Am. Chem. Soc., 65, 702 (1943).
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ISOTOPIC EXCHANGE STUDIES BY ELECTRODEPOSITION current density curves obtained by this means are shown in Figure 1. At low current densities, there is a comparatively high concentration of ionic chloride in the layer adjacent to the anode surface. This concentration may result in the formation of negative silver chloride complexes, (Ag,CI,L+,)"-, in solution. At low current densities, the formation of these complexes may compete with the electrode reaction Ag C1- 3 AgCl eand this may explain the low yields seen to the left in Figure 1 (cf. ref 16 and 17). Test of the Kinetic Procedure. According to eq 1, a semilog plot, of 1 - F vs. t should yield a straight line through the origin. A graph of this kind, shown in Figure 2, is a convenient check of the experimental procedure.'* 'The point at zero time of exchange was obtained by making a deposition in a pure lithium radiochloride solution before adding the organic chloride. Mechanism of Exchange. For a mixed SNl-sN2 reaction, in which both ions (Cl-) and ion pairs (LiC1) are involved, the rate of exchange may be represented by the equation
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R
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=: k l ~ kiaac
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+ km(1 -
~ ) U C
(2)
where kl is a first-order rate constant, ki and IC, are second-order rate constants for the reactions of ions and ion pairs, respectively, and a! is the degree of dissociation of the salt. For low salt concentrations, when k,(l - a)
