摘要:

J. Chem. SOC., Faraday Trans. 1, 1982, 78, 1065-1069 Temperature Dependence for the Reduction of Water to Hydrogen by Reduced Methyl Viologen on Platinum BY MILICA T. NENADOVIC,* OLGA I. M I ~ I ~ AND RADOSLAV R. A D ~ I C ~ Boris KidriE Institute of Nuclear Sciences, 11001 Belgrade, P.O. Box 522, Yugoslavia Received 9th April, 198 1 The electron-transfer reactions from methyl viologen monocation radicals to colloidal platinum hydrosols were studied using the pulse-radiolysis technique. Besides the reduction of water, hydrogenation of methyl viologen on platinum also takes place in these reactions. Evolution of hydrogen is enhanced at higher temperatures (65 "C) and reduction of MV+ is hindered. The reaction shows complex behaviour in the presence of hydrogen because of the formation of reduction products which decrease the catalytic activity of platinum.The reaction rate decreases from 1.3 x lo3 s-' in an argon-saturated solution to 0.6 s-' in a hydrogen-saturated solution in the presence of 2.5 x mol dm-3 Pt hydrosols. The photoinduced reduction of methyl viologen (NN'-dimethyl-4,4'-bipyridine dication, MV2+) by a suitable sensitizer can be employed to achieve the reduction of water to hydrogen over redox catalysts.' The resulting one-electron acceptor couple MV2+/MV+ has a low potential E, = -0.44 V, sufficient to reduce water: MV+ + 2H20 + iH2 + MV2+ +OH-. cat However, it has been shown recently that MV2+ irreversibly decomposes to some extent during irradiation in the presence of catalyst2? due to a hydrogenation process on platinum, which leads to a completely reduced bipiperidine derivative :3 H* MV2+ (or MV+) __+ reduction product.Pt In this work pulse radiolysis has been used to study the kinetic behaviour of reaction (1). This reaction was observed both at room temperature and 65 OC in order to determine the optimum conditions for hydrogen evolution. EXPERIMENTAL All the reagents used were commercial products of highest purity available. Solutions were prepared with triply distilled water. Oxygen was removed by bubbling with argon. The pH was adjusted with H,SO,, NaOH or phosphate buffers. Colloidal platinum was prepared as follows :, 0.02 g hexachloroplatinic acid was dissolved in 142.5 cm3 distilled water and the solution brought to the boiling point. 7.5 cm3 of 1 % sodium citrate solution was then added and the solution boiled for 1 h, water being added to supply the quantity lost by evaporation.After approximately 15 min the yellow colour began to deepen and after 30 min turned to deep brown. The solution was then stirred with a ion-exchange resin to remove the citrate. The liquid contained small particles of colloidal platinum (average size t Permanent address : Institute of Electrochemistry, ICTM and Centre for Multidisciplinary Studies, University of Belgrade, Belgrade, Yugoslavia. 10651066 REDUCTION OF WATER BY REDUCED METHYL VIOLOGEN 3 nm, measured by electron microscopy). EDTA was determined by reduction with arsenious acid.5 For determination of citric acid, an improved pyridine-acetic anhydride method was used.6 For pulse radiolysis a Febetron 707 (Field Emission Corp.) electron accelerator was used and the operation conditions were similar to those in the previous work.’ The total light path through the cell was 5.1 cm.The absorbed doses were in the range of 10-1 5 Gy per pulse. The measurements were taken at 19 f 1 O C . The gas-chromatographic method was employed for hydrogen detection using a Perkin-Elmer 154D instrument with a column of silica gel at 50 OC. RESULTS AND DISCUSSION FORMATION OF MVf IONS The reduced form of methyl viologen was produced by the reaction of hydrated electrons with MV2+ during pulse radiolysis H,O +w+ eLq, H, OH, H,O,, H,, H30+, OH- (3) eiq + MV2+ + MV+ + H,O k = 8.4 x 1Olo dm3 mol-1 s-l [ref. (S)]. (4) The nature and absorption spectrum of stable radical MV+ (Amax = 605 nm; Esos = 8.6 x lo3 mol-l cm-l) have been established previously.*.EDTA (ethylenediamine tetra-acetic acid) ( lo-, mol dm-3) was added to the solution (lov4 mol dm-3 MV2+, pH 5, unbuffered) to scavenge OH and H radicals. In the radiolysis of amino acids it is known that the a-carbon atom to the -COOH group is prone to attack by OH or H radical leading to the abstraction of hydrogenlo ( 5 ) (6) OH + RR’NCH,COOH + H,O + RR’NCHCOOH H + RR’NCH,COOH + H, + RR’NCHCOOH k = 1.2 x lo9 dm3 mol-l s-l, pH 7 [ref. (1 l)] k = 6.5 x lo7 dm3 mol-l s-l, pH 1 [ref. (12)]. It is assumed that a similar set of radicals is formed during photoinduced oxidation of EDTA in the presence of dye.13 The pK values of these radicals are probably very close to those of EDTA (pK, = 10.23, pK, = 6.16, pK3 = 2.67, pK4 = 1.99) since the pKvalues of carboxylic radicals seem to be the same as those of the parent molecule^.^^ We have found that the radical formed in reactions (5) and (6) reduces MV2+ to MV+.The yield for MV+ formation in aqueous EDTA solution was found to be 6.0 [G(MV+) = G(~L~)+G(H)+G(OH)]. In an irradiated solution of MV2+ (5 x rnol dm-3) and EDTA (0.1 mol dmd3) saturated with N,O (2.5 x lo-, mol dm-3), formation of MV+ was observed MV2+ + RR’NCHCOOH + MV+ + RR’N=CHCOOH (7) and k, = 1.4 x lo9 dm3 mol-1 s-l at pH 5 was obtained. All rate constants are summarized in table 1. MV+ is not stable at concentrations > 5 x lop5 mol dm-3, implying over-reduction of MV+ to dihydrobipyridyl. This reaction is known to occur with powerful chemical agents such as zinc or sodium dithionite.15 REACTION WITH PLATINUM HYDROSOLS MV+ radicals reduce water on a Pt catalyst according eqn (1).The first step is an electron transfer to platinum MV+ + Pt -P MV2+ + Pt-. (8)M. T. NENADOVIC, 0. I. MICIC AND R. R. ADZIC 1067 TABLE 1 .-RATE CONSTANTS FOR ELECTRON-TRANSFER REACTIONS IN METHYL VIOLOGEN-PLATINUM HYDROSOL AQUEOUS SOLUTION AT pH 5 platinum concentration species present reaction /mol dm-3 in solution rate constant RRNCHCOOH + MV2+ --* MV+ + RR'N=CHCOOHa - MV+ + Pt -+ Pt- + MVz+ 1 x 10-5-1 x 10-4 2.5 x 2.5 x MV++ Pb2+ -+ MVz+ + Pb+ e&+Pt -+ Pt-+H,O - 1 x 10-6-5 x 2.5 x - Ar (0.6-6.5) x lo3 s-I 1.5 x los dm3 mol-I s-l 0.6 s-l 8 x lod4 rnol dm-3 H, 5 x mol dm-s PbZ+ 1 x 10+ mol Pbz+, pH 5, Ar 8 x 104 mol dm-3 H, + 8 x mol dm-3 H, 3 x 103 s-1 < 10 dm3 mol-l s-' 7 x lo* dm3 mol-I s-' < los dma mol-l s-' Ar Radicals RRNCHCOOH form in reaction (5).I 1 I (a) 0 / O ' I I I I I 2 4 6 8 10 [PtJ/10-5 rnol dm-3 T/"C FIG. 1.-Rate constant for reaction of MV+ with Pt hydrosols: (a) effect of the Pt hydrosol concentration on the decay of MV+ absorption at 600 nm; (b) effect of temperature on reaction rate.1068 REDUCTION OF WATER BY REDUCED METHYL VIOLOGEN We have used colloidal platinum hydrosols stabilized with citrate ions4 This catalyst shows better reproducibility and smaller activity loss than Pt-poly(viny1 alcohol). The decay rate of MV+ by reaction (8) does not depend on the initial MV+ concentration (3 x 10-'j-5 x mol dm-3) but increases sharply with Pt concentra- tion, fig. 1 (a).The same behaviour has been noticed earlieP for centrifuged colloidal Pt stabilized by poly(viny1 alcohol). A platinum-electron adduct formed in reaction (8) is the same as the intermediate for the reaction of hydrated electrons with Pt particles (9) The rate constant for reaction (9) is 7 x lo9 dm3 mol-l s-l. We were not able to identify the transient absorption spectrum of Pt-. In the experiments where EDTA was added to scavenge OH radicals, and after the recombination of EDTA radicals (k = 1.5 x lo9 dm3 mol-1 s-l), only very weak absorption below 500 nm was obtained with a lifetime of ca. 1 s. We assumed that this absorption is probably derived from intermediates on the platinum. However, since the transient absorption is very weak, the procedure used in this study does not permit unequivocal identification of the intermediates and observation of their kinetics.The rate of decay of MV+ by reaction (8) increases with increasing solution temperature [fig. 1 (b)]. From the temperature dependence, the activation energy for reaction (8) was calculated to be 12.5 kJ mol-I, which corresponds to the diffusion- limited reaction. We have found that the hydrogen evolution in reaction (1) is enhanced at higher temperature. Hydrogen formationincreases by a factor two withincreasing temperature from 20 to 70 O C . However, even at 70 O C only 60% of MV+ ions produce hydrogen according to the stoichiometric relation of reaction (1). The rest probably disappears in the hydrogenation process on platinum [eqn (2)].eLq + Pt + Pt- + H20. HYDROGEN CONCENTRATION EFFECT In the presence of hydrogen, reaction (8) decreases markedly. The increase of hydrogen concentration decreases the rate constant value. The rate of MV+ decay at a concentration of 2.5 x loh5 mol dm-3 Pt in a hydrogen-saturated solution was found to be 0.6 s-l. Temperature appears to affect remarkably the rate of this reaction and the increase of the rate constant is not only a result of the temperature effect on the reaction activation, but also due to the nature of platinum surface in the presence of hydrogen. The inhibition of electron transfer to Pt most likely reflects the formation of some products which are strong poisons for Pt or effect the coagulation of Pt particles. MV+ and MVO are produced upon saturating an MV2+ and Pt colloid containing solution with H,: (10) (1 1) We were able to detect a MV+ steady-state concentration of 3 x mol dm-3 by measuring the absorbance at 605 nm; the hydrogenation process of MV2+ proceeds via formation of MV+ ions in a solution of 5 x mol dm-3 MV2+, 1 x mol dm-3 EDTA and 2.5 x lo-* mol dm-3 Pt saturated with hydrogen.This hydrogenation process leads to complete destruction of MV2+ ions. On the other hand, the formation of MV+ and MVO decreases the catalytic activity of the platinum. Coagulation of Pt hydrosols can be observed when a colloidal solution of Pt containing MV2+ stands for more then 30 min after saturation with H,. MV2+ + iH2 + MV+ + H+ MV+ + $H2 + MVO + H+.M. T. NENADOVIC, 0.I. MICIC AND R. R. ADZIC 1069 On principle, the presence in the solutions of any species more strongly adsorbed than hydrogen will change the mechanism of reactions (1) and (2). We chose Pb2+ ions since they adsorb strongly on the Pt surface17 and modify the catalytic properties of platinum.18 MV+ ions do not react with Pb2+ when Pt hydrosols are absent (table 1). However, we found that Pb2+ ions affect electron transfer from MV+ to Pt. The presence of 5 x mol dm-3Pb2+ ions increases the reaction rate from 0.6 to 3 x lo3 s-l in a hydrogen-saturated solution. Though the second step, hydrogen evolution, is inhibited when Pb is adsorbed on platinum,18 our experiments show that reduction of water on Pt by MV+ ions is a complex process sensitive to many factors.Electron transfer from MV+ to Pt in the presence of hydrogen is a real condition which can be expected in the case of continuous evolution of hydrogen from water. Note that under this condition, undesirable reactions, such as the hydrogenation of MV2+, take place and drastically decrease the catalytic activity of platinum. B. V. Koryakin, T. S. Dzhabiev and A. E. Shilov, Dokl. Akad. Nauk SSSR, 1977, 238, 620; K. Kalyanasundaram, J. Kiwi and M. Gratzel, Helv. Chim. Acta, 1978, 61, 2720; A. Moradapour, E. Amouyal, P. Keller and H. Kagan, Nouv. J. Chim., 1978, 2, 542; P. J. De Laive, B. P. Sullivan, T. J. Meyer and D. G. Whitten, J. Am. Chem. Soc., 1979, 101, 4007; T. Kawali, K. Tanimura and T. Sakada, Chem. Lett., 1979, 137; M. Kirsch, J. M. Lehn and J. P. Sauvage, Helv.Chim. Acta, 1979, 62,1345; K. Kalynasundaram and M. Gratzel, J. Chem. Soc., Chem. Commun., 1979,1138; 0. I. Midid and M. T. Nenadovid, J. Chem. Soc., Faraday Trans. 1, 1981, 77, 919. M. Gohn and N. Getoff, Z. Naturforsch., Teil A , 1979, 34, 1135. P. Keller and A. Moradpour, J. Am. Chem. Soc., 1980, 102, 7193. G. C. Bond, Trans. Faraday Soc., 1956, 52, 1235. 26, 1806. Y. R. Maner and M. Boulet, Anal. Abstr., 1959, 6, 3762. ti P. V. Gherney, B. Crafts, H. H. Hagermoser, A. Y. Boule, R. Harbin and B. Zak, Anal. Chem., 1954, ' V. M. Markovid, D. Nikolid and 0. I. Midid, Int. J. Radiat. Phys. Chem., 1974, 6, 227. * J. A. Farrington, M. Ebert, E. J. Land and K. Fletcher, Biochim. Biohys. Acta, 1973, 314, 372. lo W. M. Garrison, Radiat Res. Suppl., 1964, 4, 158. l1 Farhataziz and A. B. Ross, Selected SpeciJic Rates of Reactions of Transients from Water in Aqueous Solution (NSRDS-NBS 59, Washington D.C., 1973). l2 M. Anbar, Farhataziz and A. B. Ross, Selected Specific Rates of Reactions of Transients from Water in Aqueous Solution (NSRDS-NBS 51, Washington D.C., 1973). l3 P. Keller, A. Moradpour, E. Amouyal and H. Kagan, Nouu. J, Chim., 1980,4, 377. l4 P. Neta, M. Simic and E. Hayon, J. Phys. Chem., 1969, 73, 4207. l5 I. G. Carey, J. F. Cairns and J. E. Colchester, J. Chem. Soc. D, 1969, 1280. l6 J. Kiwi and M. Gratzel, J. Am. Chem. Soc., 1979, 101, 7214. l7 D. M. Kolb, in Advances in Electrochemistry and Electrochemical Engineering, ed. H. Geriseler and A. I. Krasna, Photochem. Photobiol., 1979, 29, 267. C. Tobias (J. Wiley, Chichester, 1978), vol. 11. R. R. Adid, Isr. J. Chem., 1979, 18, 166. (PAPER 1/567)

ISSN:0300-9599    

DOI:10.1039/F19827801065

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1982, 78, 1071-1077 Thermodynamic Properties of Copper Sulphate in Dioxan + Water Mixtures from Electromotive-force Measurements BY RATAN L. BLOKHRA* AND SUDARSHAN KOHLI Department of Chemistry, Himachal Pradesh University, Simla- 17 1005, India Received 10th April, 1981 Electromotive-force measurements of the cell Cu-HglCuSO, (m), dioxan (X)IHg,SO,(s)lHg have been made at 288, 298, 308 and 318 K for solvent compositions X = 10, 20, 30 and 40% (w/w) of dioxan. These have been used to evaluate the standard potentials of the cell, the mean activity coefficient of copper sulphate (molal scale) and the thermodynamic functions of transfer of copper sulphate from water to the respective dioxan + water media. The results have been interpreted in terms of solute-solvent interactions.The activity coefficients of many typical 1 : 1, 2 : 1 and 2 : 2 ele~trolytesl-~ have been determined in various aquo-organic solvents and a few non-aqueous solvents. The thermodynamics of aqueous solutions of cadmium, copper, manganese and nickel sulphates have not been studied as comprehensively as those of aqueous solutions of zinc ~ u l p h a t e . ~ ? ~ A survey of the literature reveals that very little work has been reported on thermodynamic studies of 2: 2 electrolytes in aquo-organic and non- aqueous solvents. Recently Blokhra et ~ 1 . ~ 9 lo have reported on the thermodynamics of copper sulphate in diethylene glycol, aqueous diethylene glycol, ethylene glycol and aqueous ethylene glycol. The present investigations have been carried out using copper sulphate in order to determine (i) the activity coefficient of copper sulphate in various aqueous dioxan mixtures and (ii) the thermodynamic functions of transfer of copper sulphate from water to the respective dioxan +water media.The solvent systems chosen were 10, 20, 30 and 40% (w/w) dioxan+ water mixtures. EXPERIMENTAL Dioxan was of B. D. H. AnalaR quality and was purified by refluxing over sodium for 6 h followed by distillation. Water of specific conductance of ca. kg mol-l i2-l was used for making up, by weight, aqueous mixtures of dioxan. Copper sulphate (AX.) was used as such. Experimental solutions of the desired concentrations were prepared by weight. The e.m.f. cell used for the present study was as previously describedg except that a copper amalgam was used instead of the electroplated copper electrode in one of the limbs.Copper amalgam containing ca. 3 % copper (w/w) was prepared by electrolysing a copper(I1) perchlorate solution with a mercury pool as cathode, as described e1sewhere.l' At room temperature, a two-phase amalgam is formed between 0.0032 and 24.1 % copper (w/w).l29 l3 Within this range the potential will be independent of composition, and the exact composition of copper is therefore unimportant.14 A high copper content should be avoided, however, as the amalgam then becomes inconveniently stiff. The amalgam was stored under 0.1 mol dm-3 perchloric acid. Prior to use it was washed with dilute perchloric acid to remove traces of copper(1r). The e.m.f. measurements were made with an OSAW (Ambala) precision potentiometer 10711072 PROPERTIES OF CUSO, IN DIOXAN AND WATER having an accuracy of kO.1 mV.A d.c. spot-galvanometer was used in conjunction with the potentiometer. The potentiometer was standardised against a certified Weston standard cell maintained at a constant temperature. All measurements were made in a thermostat having fluctuations of < k0.2'. The cell attained equilibrium after 30-35 min in all the solvent mixtures. The concentrations of the solutions were occassionally checked following the experiments and no significant change was detected. Duplicate experiments were performed simultaneously in each case and the duplicates generally agreed within kO.5 mV. The densities (Po), dielectric constant (D,) and the Debye- Hiickel parameters for the various dioxan+water mixtures were taken from the work of Das et aL2 RESULTS AND DISCUSSION The e.m.f.(E) values of the cell are reported in table 1. Assuming copper sulphate to be fully dissociated, the standard potentials, E g , of the cell: Cu-HglCuSO, (m), dioxan (X)IHg,SO, (s)(Hg in different solutions can be estimated from an equation of the Hitchcock15 type: where I is the ionic strength of the solution, m is the molal concentration of the solution, A and B are the Debye-Huckel constants, a, is the ion-size parameter, Z+ and 2- are the valencies of the cation and anion, respectively, and k equals 2.3026 RTIF. B' is a constant quantity and it is a measurement of the interaction energy.16 The values of the activity coefficients, A,, - can be evaluated from the relation17 E = @ - k log mA+.- (2) should be a linear function of I when a suitable value of the ion-size parameter is chosen. The intercept of the plot at I = 0 then gives the value of E,e. The a, values for the various solvent mixtures at different temperatures are chosen in such a way that the deviation from linearity of the plot of Ei against I is minimum. A deviation of k 0.2 mV was observed when the a, value was vaned within k0.3 A of the chosen value. The following a, values were selected for the respective dioxan + water mixtures : From eqn ( 1 ) it is expected that dioxan (wt %) 10 20 30 40 aolA 5.0 5.2 5.2 5.5 a, is constant throughout the temperature range studied for a particular solvent composition.This agrees with the conclusions of LaMer et aZ.6*s The values of are recorded in table 2. The average standard deviation in is kO.2 mV and the values for each solvent system were fitted by the method of least squares to eqn (3) (3) where T is the temperature in K. The constants a, b and c are given in table 3. The mean activity coefficient A, of copper sulphate in various solvent media was calculated with the help of eqn (2), and these values at 298 K are listed in table 4. An error of kO.05 mV in the e.m.f. values results in an error of kO.001 in the value of 1,. - The values of the activity coefficient at a particular molality are found to = a- b(T- 298.15) - C( T - 298.R. L. BLOKHRA A N D S. KOHL1 1073 TABLE 1 .-ELECTROMOTIVE-FORCE MEASUREMENTS OF THE CELL USED (E/V) IN VARIOUS DIOXAN +WATER MIXTURES AT DIFFERENT TEMPERATURES T / K concentration/ mol kg-l 288 298 308 318 1 .oo 2.00 4.00 6.00 8.00 10.00 20.00 40.00 60.00 80.00 100.00 1 .oo 2.00 4.00 6.00 8.00 10.00 20.00 40.00 60.00 80.00 100.00 1 .oo 2.00 4.00 6.00 8.00 10.00 20.00 40.00 60.00 80.00 100.00 1 .oo 2.00 4.00 6.00 8.00 10.00 20.00 40.00 60.00 80.00 100.00 0.5204 0.5062 0.4929 0.4849 0.4799 0.4763 0.4652 0.4553 0.4506 0.4475 0.4452 0.5152 0.5016 0.488 1 0.48 14 0.4762 0.4723 0.4617 0.4522 0.4470 0.4434 0.4405 0.5070 0.4933 0.48 17 0.4741 0.4699 0.4658 0.4554 0.4450 0.4376 0.4319 0.4269 0.5005 0.4880 0.4769 0.4707 0.4664 0.4632 0.4541 0.4438 0.4358 0.4290 0.4228 10% dioxan 0.5227 0.5077 0.4940 0.4864 0.48 1 3 0.4772 0.4660 0.4560 0.451 1 0.4483 0.4467 20% dioxan 0.5170 0.5029 0.4893 0.482 1 0.4773 0.4739 0.4633 0.4536 0.4484 0.445 1 0.4426 30% dioxan 0.5077 0.4939 0.48 15 0.4748 0.4702 0.4666 0.4566 0.4472 0.4417 0.4377 0.4345 0.5029 0.4902 0.4788 0.4726 0.4682 0.4649 0.4564 0.4464 0.4392 0.4334 0.4282 40% dioxan 0.5247 0.5094 0.4954 0.4877 0.4826 0.4787 0.4670 0.4563 0.45 17 0.4490 0.4473 0.5185 0.5040 0.4907 0.4836 0.4783 0.4748 0.4639 0.4543 0.4493 0.446 1 0.4439 0.5099 0.4960 0.4833 0.4768 0.47 18 0.468 1 0.4579 0.4485 0.4429 0.4388 0.4357 0.5048 0.49 17 0.4803 0.4742 0.4700 0.4668 0.4582 0.4482 0.4419 0.4365 0.43 19 0.5266 0.51 13 0.4968 0.4889 0.4836 0.4795 0.4680 0.4573 0.4537 0.45 12 0.4500 0.5198 0.5046 0.49 14 0.4838 0.4786 0.475 1 0.4642 0.4555 0.4503 0.4473 0.4455 0.51 14 0.4970 0.4846 0.4775 0.4726 0.4688 0.4595 0.4498 0.4442 0.4402 0.4368 0.5064 0.4933 0.48 17 0.4757 0.4714 0.4683 0.4594 0.4504 0.4444 0.4394 0.43481074 PROPERTIES OF CUSO, I N DIOXAN AND WATER decrease with increasing dioxan content in the medium, as expected from Debye- Huckel theory, because of a lowering of the dielectric constant of the medium.The higher magnitude of the activity coefficient at lower dioxan contents suggests that the solute-solvent interaction increases as the dioxan content decreases. It is also seen that the activity coefficient at a particular molality decreases with increasing temperature. This may be attributed to a greater solute-solvent interaction at lower temperatures. TABLE 2.-sTANDARD MOLAL POTENTIALS OF THE CELL USED IN DIOXAN + WATER MIXTURES AT VARIOUS TEMPERATURES dioxan (wt %) 288 K 298 K 308 K 318 K ~~ ~ 0 0.3455 0.3420 0.3385 0.3350 10 0.3410 0.3370 0.3325 0.3280 20 0.3340 0.3295 0.3245 0.3190 30 0.3235 0.3 175 0.3130 0.3075 40 0.3140 0.3095 0.3040 0.2975 TABLE VALUES OF THE CONSTANTS a, b AND c IN EQN (3) b/ 10-4 v K T ~ c/ V K-2 dioxan (wt %) a/v 0 0.3420 10 0.3372 20 0.3294 30 0.3179 40 0.3095 3.00 4.25 4.72 5.25 - 4.98 2.50 1.70 2.40 0.20 2.60 The free-energy changes AGP accompanying the transfer of one mole of copper sulphate from the standard state in water to the mixed solvents were calculated on the mole-fraction scale using the relationls AGP = - 2F [(Efit), - (Efit),] - 2 x 2.3026 RT log (MJM,).(4) The corresponding entropy changes were calculated using the relation1* A@ = F(b, - b,) + 2F( T - 298.15) (c, - c,) + 2 x 2.3026 R log (M,/M,) ( 5 ) where the subscripts w and s refer to water and the mixed solvent, respectively.The corresponding enthalpy changes will be given by the equation AH? = AG?+ T A F . (6) The AGP, AH? and ASP values are listed in table 5. The probable uncertainties in AGP are f 19 J mol-l, those in AH? are &23 J mol-l, and those in A@ are k0.40 J K-l mol-l in 10 and 20% solvent composition and kO.70 J K-l mol-1 in 30 and 40% solvent composition. The standard Gibbs free energies of transfer, AGP, are positive for all the solvent compositions and increase with increasing temperature. The positive AGP values suggest that copper sulphate is in a higher free-energy state in dioxan +water mixtures than in water so that copper sulphate hasmore affinity for water than for dioxan + waterR.L. BLOKHRA AND S. KOHL1 1075 mixtures, and the transfer of copper sulphate from water to dioxan +water mixtures is not a spontaneous process with the solute in the standard state in either medium. Similar conclusions were drawn for copper sulphate in ethyleneglycol + water mixtures.l0 The values of A p and AH? are also positive for all the solvent mixtures. The enthalpy in dioxan+ water mixtures is, therefore, greater than that in pure water and hence the increase in order created by copper sulphate in dioxan + water mixtures is less than that occurring in pure water. It may be possible to split the AGP values TABLE 4.-MEAN MOLAL ACTIVITY COEFFICIENT OF CUSO, IN VARIOUS DIOXAN +WATER MIXTURES AT 298 K ~~ ~ dioxan (wt %) concentration/ mol kg-l 10 20 30 40 1 .oo 2.00 4.00 6.00 8.00 10.00 20.00 40.00 60.00 80.00 100.00 0.725 0.65 1 0.555 0.498 0.455 0.426 0.330 0.243 0.196 0.164 0.140 0.676 0.585 0.497 0.438 0.396 0.362 0.274 0.199 0.163 0.139 0.122 0.609 0.522 0.423 0.366 0.328 0.302 0.223 0.160 0.132 0.116 0.105 0.538 0.441 0.344 0.291 0.259 0.236 0.165 0.121 0.107 0.100 0.098 TABLE 5.-FREE ENERGY, ENTHALPY AND ENTROPY OF TRANSFER OF COPPER SULPHATE FROM WATER TO DIOXAN + WATER MIXTURES AT VAROUS TEMPERATURES T / K AGP/J mol-l AHp/kJ mol-l A P / J K-l mol-l 288 298 308 318 288 298 308 '318 288 298 308 318 288 298 308 318 469 55 1 73 1 910 1389 1552 1813 2170 2936 3373 3521 3861 424 1 43 70 4692 5207 10% dioxan 4.78 4.56 4.40 4.20 7.05 7.35 7.75 7.86 12.00 11.19 10.00 8.89 11.52 11.96 12.60 13.43 20% dioxan 30% dioxan 40% dioxan 14.98 13.44 1 1.90 10.35 19.66 19.47 19.28 19.08 3 I .46 26.25 21.04 15.83 25.29 25.48 25.67 25.871076 PROPERTIES OF CUSO, IN DIOXAN A N D WATER into two parts, as suggested by Roy et al.,19 a non-electrostatic or chemical contribution AG?,, and an electrostatic contribution AG?,, which has been calculated from the Born equation,20 eqn (7): where N is Avogadro's number and D, and D, are the dielectric constants of the mixed solvent and water, respectively; r+ and r- are the radii of the Cu2+ and SO:- ions, taken as 0.70 and 2.89 respectively.TABLE 6.-ELECTRICAL AND THE CHEMICAL PARTS OF THE THERMODYNAMIC QUANTITIES ACCOM- PANYING THE TRANSFER OF COPPER SULPHATE FROM WATER TO DIOXAN -I- WATER MIXTURES AT 298 K dioxan AG?el! AGfch/ A H r e l / AHrch/ A q e , / A q c h / (wt %) kJ mo1-1 kJ m o P kJ mol-1 kJ mol-l J K-' mol-l J K-l mol-l ~~ 10 1.92 - 1.37 - 2.8 1 7.37 - 15.87 29.3 1 20 4.53 - 2.98 - 5.34 12.69 - 33.10 52.57 30 8.00 - 4.63 - 8.50 19.69 - 55.38 8 1.63 40 12.92 -8.55 -12.93 24.89 - 86.77 1 12.25 The electrostatic part of the entropy of transfer may be obtained by differentiating eqn (7), whereby we have (8) where the values of d In D,/dT and d In D,/dT can be evaluated from the simple empirical eqn (9):2 dlnD 1 d T 8 in which 8 is a constant characteristic of the medium.Thus eqn (8) may be rewritten as 1 dlnD, - -- ----- - ~ ( ~ w d i ~ ~ w D, dT )(;+:) (9) -- -- - From the slopes of the linear plots of log D against T for the respective dioxan +water mixtures, the following values of 8 were calculated: dioxan (wt %) 0 10 20 30 40 8 220 202 194 187 181 From a knowledge of AGee1 and A T , , the electrostatic part of the enthalpy change AH?,l has been computed.The chemical contribution of the free energy of transfer AGFCh, entropy of transfer A T , , and enthalpy of transfer can then be obtained by subtracting the respective electrostatic contribution values from the molal quantities. These values so calculated at 298 K are presented in table 6. It is obvious that the chemical contribution of the free energy of transfer is negative and appearsR. L. BLOKHRA AND S. KOHLI 1077 to be a parameter of the solvent which measures the increase in basicity in the dioxan +water mixture. Hence, considering only the chemical contribution to the free energy AGP which has negative values, the dioxan +water mixture appears to be more basic than water.The electrostatic factors, however, predominate over the chemical contribution or the solvation, resulting in an overall unfavourable effect on the transfer process from water to dioxan + water mixtures. The electrostatic parts of the enthalpy and entropy have negative values, whereas the chemical contribution to the enthalpy and entropy is positive. We thank the U.G.C. (India) for financial support. S. K. thanks the D.A.V. Colleges Managing Committee, New Delhi, for sanctioning leave under the faculty improvement programme. P. K. Das and U. C. Misra, Electrochim. Acta, 1977, 22, 59. B. K. Das and P. K. Das, J.Chem. Soc., Faraa'ay Trans. 1, 1978, 74, 22. K. K. Kundu and K. Mazumdar, J. Chem. Soc., Faraday Trans. I , 1975, 71, 1422. R. F. Newton and E. A. Tippets, J. Am. Chem. SOC., 1936,58, 280. H. S. Harned and J. C. Hecker, J. Am. Chem. Soc., 1934, 56, 650. V. K. LaMer and W. G. Parks, J. Am. Chem. Soc., 1931, 53, 2040. I. A. Cowperthwaite and V. K. LaMer, J. Am. Chem. Soc., 1931, 53, 4333. R. L. Blokhra, Y. P. Sehgal and V. K. Kuthiala, Electrochim. Acta, 1976, 21, 1079. 'I U. B. Bray, J. Am. Chem. Soc., 1927,49, 2372. lo R. L. Blokhra and Y. P. Sehgal, Znd. J. Chem., 1977, 15(A), 1035. l1 S. Ahrland and B. Tagesson, Acta Chem. Scand., Ser. A, 1977, 31, 615. l2 G. Tammann and J. Z . Stassfurth, Znorg. Chem., 1927, 160, 246. l 3 G. Tammann and K. Z. Kollmann, Znorg. Chem., 1925, 143, 357. l4 S. Ahrland and J. Rawsthome, Acta Chem. Scand., 1970, 24, 157. l5 D. 1. Hitchcock, J. Am. Chem. Soc., 1928, 20, 2076. l6 D. J. G. Ives and G. J. Janz, Reference Electrodes (Academic Press, New York, 1961), p. 38. l7 D. A. MacInnes, The Principles of Efectrochemistry (Dover Publications, New York, 1967), p. 165. lS R. N. Roy, W. Vernon and A. L. M. Bothwell, Electrochim. Acta, 1972, 17, 5. 2o M. Born, 2. Phys., 1920, 1, 45. *l L. Pauling, The Nature of the Chemical Bond (Cornell University Press, Ithaca, 3rd edn, 1960), p. 521. K. Bose, A. K. Das and K. K. Kundu, J. Chem. Soc., Faraa'ay Trans. I , 1975, 71, 1838. (PAPER 1/581)

ISSN:0300-9599    

DOI:10.1039/F19827801071

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. Soc., Faraday Trans. 1, 1982, 78, 1079-1086 Enantiomer-differentiating Hydrogen Transfer from Fenchol to Menthone over Molten Indium Catalyst B Y HIROSHI SUGAWARA AND YOSHISADA OGINO* Department of Chemical Engineering, Faculty of Engineering, Tohoku University, Aramaki Aoba, Sendai 980, Japan Received 13th April, 1981 The initial rate (To) of the hydrogen-transfer reaction from D-fenchol to menthone has been found to obey the following equation where k is the rate constant, p A is the partial pressure of D-fenchol, and the subscript j denotes any one of the optical isomers (L-,D- and D,L-) of menthone. The reaction is enantiomer-differentiating: kL > kD, and the optical resolution of D,L-menthone takes place in the reaction between D-fenchol and Dpmenthone. The experimental results are explained on the bases of an Eley-Rideal-type mechanism and stereo- chemically consistent transition-state models.Probable adsorption models for menthone are also proposed. '0.j = k* P A Heterogeneous enantiomer-differentiating reactions are closely related to bio- chemical enzymic reactions;l thus an unambiguous understanding of the mechanisms of the former reactions provides us with strong bases for studying those of the latter. A heterogeneous enantiomer-differentiating reaction is usually studied using a batch reaction system consisting of gas, liquid and solid (catalyst) phases.2v This introduces many difficulties in analysing the reaction rate, and therefore in clarifying the reaction mechanism: one must clarify the effects of agitation and mass transfer, effects of solvent and of solution pH, besides effects caused by the choice of catalyst used.Recently, we4y5 have found that liquid metals such as liquid indium are able to catalyse certain kinds of vapour-phase enantiomer-differentiating reactions. Since kinetic formalism as a tool for investigating the mechanism of a vapour-phase heterogeneous catalytic reaction is firmly established,s this finding shows the potential merit in clarifying the mechanism of enantiomer differentiation. In the present work, the kinetics of the enantiomer-differentiating hydrogen transfer from fenchol to menthone over a liquid-indium catalyst has been studied in order to discuss the reaction mechanism. EXPERIMENTAL MATERIALS Special attention was paid to the purity of the reagents used.The purity of D-fenchol (the hydrogen-donating reagent) was better than 98.5% and that of cumene (the solvent for the reactants) was better than 98 %. Both L-menthone and D,L-menthone (hydrogen-accepting reagents) were purified by means of column chromatography (silica gel, Wako-C-200, 100 g; benzene eluent) followed by distillation to remove benzene. The purity of the resulting menthone was better than 98%. The purity of indium metal (the catalyst) was 99.999%. The catalyst metal was fused, reduced and purified by a previously reported method.' 10791080 H, TRANSFER FROM FENCHOL TO METHONE APPARATUS The reaction apparatus used in the present work is shown schematically in fig. 1. This is essentially a conventional flow-type apparatus except for the reactor, which is of the rectangular- duct type.The use of a reactor of this type greatly facilitates kinetic measurements with liquid-metal catalysts.* A cumene solution of the required concentrations of D-fenchol and L-menthone (or D,L- menthone) was supplied from an electrically driven micro-feeder into the reaction system, Within the preheating zone the feed solution was evaporated and mixed with a stream of purified helium. The mixture was then led to the reactor and subjected catalysis at the liquid-indium surface. The effluent from the reactor was cooled and the liquid products were separated from helium and the gaseous products (mainly hydrogen). I F I I I I I Gl ' J ! I G3 '- - 8 FIG. 1 .-Reaction apparatus with a rectangularduct-type reactor: (A) microfeeder, (B) preheater, (C,), (C,) electric furnaces, (D) rectangular-duct-type reactor (86 mm length, 20 mm width, 16 mm height), (E) cooler, (F) separator, (G,), (G.J, (G3) Dewar vessels, (H) sampling port, (I) soap-film flow meter, (J) trap containing molecular sieves, (K) vessel containing reduced copper catalyst, (L) helium reservoir.KINETIC STUDIES Kinetic data were obtained by keeping the catalyst surface area (S) constant (ca. 18 cm2) and varying the feed rate (F) of the reactants: the conversion (1,) for the hydrogen-transfer reaction was measured as a function of the contact time ( S / F ) . Thus the initial reaction rate (r,,) was determined from the initial slope of the curve representing the relation between X, and S/F. Independent variables adopted in the rate study were the reaction temperature (T) and the respective partial pressures of fenchol (pK) and menthone (PA).ANALYSIS Analysis of the liquid products was carried out mainly by gas chromatography with an apparatus (Hitachi model 163) equipped with a capillary column (Unicon oil, 0.25 mm x 45 m) and a flame ionization detector. Each compound produced was identified by mass spectrometry (JEOL-JMS-D-300 gas-chromatographic mass spectrometer). In order to clarify if the optical resolution of D,L-menthone resulted from hydrogen transfer from D-fenchol to D,L-menthone, the specific optical rotation of the unreacted menthone remaining in the reaction products wasH. SUGAWARA AND Y. OGINO 1081 measured at 15 "C with an automatic digital polarimeter (Union Giken PM-101). For this measurement menthone was separated from the products by the following method; first cumene (solvent) was removed by distilling the liquid products, and then menthone was separated from the residual mixture by means of column chromatography (silica gel column; benzene eluent).Finally benzene (eluent) was removed by distillation. RESULTS OVERALL REACTION Product analyses revealed that (i) the liquid products were fenchone, menthol and neomenthol, (ii) the molar ratio of fenchol to (menthol + neomenthol) was always greater than unity, and (iii) the gaseous product was hydrogen. Both menthol and neomenthol could be accounted for as products of the hydrogen transfer from fenchol to menth~ne,~ and the excess of fenchone and formation of hydrogen could be accounted for by the dehydrogenation of fenchol.Good mass balances were obtainable with these two reactions. Furthermore, separate experiments proved that no direct hydrogenation of menthone by gaseous hydrogen took place over the liquid-indium catalyst. Thus it was considered that the overall reaction is a composite of the following reactions : fenchol + menthone - fenchone +menthol and neomenthol (1) (2) fenchol - fenchone + H,. Frequently, a side reaction prevents one from obtaining precise kinetic data for the reaction of interest. However, this was not the present case. The dehydrogenation reaction (2) led to little difficulty in obtaining the conversion (Xt) for the hydrogen- transfer reaction (1): Xt was readily obtained from the concentrations of menthol and neomenthol in the products.Thus we were able to obtain the reliable kinetic data given below. INITIAL RATE AND ENANTIOMER DIFFERENTIATION Exemplified in fig. 2 is the relation between Xt and S/F at 410 O C . This figure illustrates two important aspects of the rate of the hydrogen-transfer reaction. First, the Xt against S / F plot is linear under the experimental conditions. This greatly facilitates evaluation of the initial rate (ro), defined by [dX,/d(S/P)],,,,,. Secondly, D-fenchol reacted with L-menthone at an increased rate than with D,L-menthone. This implies that the hydrogen-transfer reaction is enantiomer-differentiating : hydrogen transfer from D-fenchol to L-menthone is faster than that to D-menthone. In contrast to this, the rate of dehydrogenation of fenchol [reaction (2)] was found to be affected little by the change of the coexisting menthone species (from L- to D,L-).The kinetic characteristics of hydrogen-transfer are best expressed in fig. 3. The reaction obeyed first-order kinetics with respect to the partial pressure of D-fenchol and zero-order kinetics with respect to the partial pressure of menthone. These characteristics were independent of the species (L- or D,L-) of menthone used as the hydrogen acceptor. However, the rate of hydrogen transfer to L-menthone was greater than that to D,L-menthone (rO,D,L). The experimental data mentioned above are summarized by the following rate equations : with1082 H, TRANSFER FROM FENCHOL TO METHONE 0 FIG.2.4onversion (X,) for hydrogen transfer as a function of contact time ( S / F ) at 410 OC,pA = 0.05 atm and pK = 0.1 atm (1 atm = 101 325 Pa): 0, D-fenchol/L-menthone system; 0, D-fenchol/D,L-menthone system. PA /atm 0.05 0.1 0 0.1 5 0.20 pK/atm FIG. 3.4haracteristic features of the initial rate (To) of the hydrogen-transfer reaction at 410 OC: 0,6, D-fenchol/L-menthone system; 0, 6, D-fenchol/D,L-menthone system; pK = 0.05 atm for rO/pA relation, and p A = 0.05 atm for rO/pK relation.H. SUGAWARA AND Y. OGINO 1083 where k is the rate constant and the subscripts L and D,L indicate that the hydrogen acceptors are L-menthone and D,L-menthone, respectively. The experimental data given in fig. 3 and the rate equations given above enabled us to evaluate the rate constant kD defined by ‘0,D = kDPA ( 5 ) where the subscript D indicates that the hydrogen acceptor is D-menthone.shown in fig. 4 reveal that demonstrating the clear enantiomer-differentiating characteristics shown by the hydrogen-transfer reaction. Thus a comparison between kL and k , becomes possible. The Arrhenius plots k, ’ k D (6) I 1 1.30 1.35 1.40 1.45 1.50 FIG. 4.-Arrhenius plots for kL (0) and kD (0). kL = 1.2 x 10gexp(-40.3 x 103/RT) k,, = 1.2 x 10gexp(-41.6 x 103/RT). The inequality (6) suggests that the optical resolution of D,L-menthone takes place in the reaction between D-fenchol and D,L-menthone. In fact, the specific rotation of menthone separated from the reaction products was + 2.08 (C = 3.072 g per 100 cm3; ethanol solvent), indicating that optical resolution took place (optical yield x 57%, the reaction temperature was 48OOC and the composition of the feed was menthone/cumene = 1.5 and cumene/fenchol = 18.6, both in molar ratios).DISCUSSION DERIVATION OF THE RATE EQUATION The reaction mechanism deserves special discussion, since liquid-indium itself has no chirality and hence the enantiomer-differentiation observed in the hydrogen-transfer reaction over this catalyst might appear unusual. Thus the present discussion mainly1084 H, TRANSFER FROM FENCHOL TO METHONE treats this problem, and our first task is to disclose the relationship between the rate equation and the underlying microscopic processes. It is possible to show that the following reaction mechanism well explains the experimental rate data.(1) The surface of the liquid-indium catalyst is almost completely covered by adsorbed menthone molecules (K) : the adsorption is saturated and the number of adsorbed K molecules (nO,) is assumed to be lo1* molecule The adsorbed K molecules are regarded as active centres with no translational or rotational freedom. (2) Chemisorption of fenchol molecules (A) upon active centres (adsorbed K molecules) limits the reaction rate. Under the assumptions mentioned above we can obtain the following rate equation :lo where k is Boltzmann’s constant, h is Planck’s constant, I is the moment of inertia of a fenchol molecule, mA is the mass of a fenchol molecule, andf, is a vibrational partition function (superscripts $, A and K denote the transition state, fenchol and menthone, respectively). Eqn (7) reduces to ro = (pA/kT) ng h4 [87r21(27rrnAkT)~]-1 (rs,/f&fl) exp (- E/RT) (7) rO = kthPA where kth is the rate constant by excluding p A from the right-hand side of eqn (7).It is clear that the theoretical rate equation is of a form identical with that of the experimental rate equations (3 a) or (3 b). We can define the theoretical frequency factor At, by k,, = At, exp ( - E/ R T). (9) With approximated values for nO,, I, f:, f: and fj, it is possible to show that the frequency factor At, can take a value comparable to the experimental value : the details of calculation are not important for the discussion of enantiomer differentiation and are not presented here. TRANS IT I ON-S TAT E MODEL ; EN ANT I 0 ME R DIFFERENT I AT I ON It is clear that the relation kL > k, expressing the degree of enantiomer differentiation observed experimentally is explainable if we can justify the relation k,,,, > kth,D. Stereochemistry tells that the relation is justified only when the following conditions are satisfied : (i) the transition state for hydrogen transfer to L-menthone is a diastereomer of the transition state for hydrogen transfer to D-menthone: (ii) the formation of the former transition state is easier than that of the latter.Whether these conditions are acceptable or not can be examined by making transition-state models. Illustrated in fig. 5 are the probable transition-state models for the following hydrogen transfer reactions : D-fenchol+ L-menthone - A 1 - L-fenchone + L-menthol D-fenchol+ D-menthone - A2 - L-fenchone + D-menthol D-ferichol+ D-menthone - B 1 - L-fenchone + D-neomenthol D-fenchol+ D-menthone - B2 - L-fenchone + L-neomenthol where Al, A2, B1 and B2 are the transition states defined in fig.5 . One basic assumption used in making these models is that the hydroxy group in the fenchol molecule combines with the carbonyl group in the menthone molecule to form the state shown in fig. 5(c). In addition it is implicitly assumed that theH. SUGAWARA A N D Y. OGINO 1085 ( A l l (B1) D - fenchol D :f enc hol (A21 ( 62) D - fenchol D - f enchpl D - ment hone FIG. 5.-Transition-state models : arrow indicates an area of large steric hindrance. formations of both menthol and neomenthol obey identical kinetic rules, while the structures of the transition states of these reaction are stereochemically different from each other.From geometries of these models, it is apparent that A1 is a diastereomer of A2, and B1 is a diastereomer of B2. Futhermore, both in the formation of A2 and in the formation of B2, points of large steric hindrance appear in the areas indicated by the arrows in fig. 5 . On the other hand, little steric hindrance appears in forming A1 and B1. Thus the relation kth,& > kth$ representing enantiomer differentiation is reasonable. In principle, it is expected that the differences in the structures of mutually corresponding pairs of diastereomers (A1 and A2; B1 and B2) should bring about differences both in the frequency factors and in the activation energies. At first sight the experimental result shown in fig. 4 seems to indicate that AL = A , and EL < ED.However, the difference in activation energies is < 8 kJ mol-l and its accuracy is uncertain, Therefore we can draw no conclusion from the experimental result. ADSORPTION MODEL Following from the reaction mechanism discussed above, the reactive menthone molecule adsorbed on the catalyst surface should be oriented according to the following restrictions: (i) the carbonyl group must be activated by contacting with the catalyst surface; (ii) the same group is open to attack by the fenchol molecule impinging from the gas phase. Illustrated in fig. 6 are the probable adsorption models obtained by taking these1086 H, TRANSFER FROM FENCHOL TO METHONE FIG. 6.-Possible models for the reactive L-menthone molecule adsorbed on the liquid-indium catalyst : arrow indicates the direction from which the fenchol molecule attackg: 0, hydrogen; 0, carbon; 0, oxygen.restrictions into consideration. The models are for adsorbed L-menthone : the mirror image of each gives the corresponding adsorbed state for D-menthone. From the adsorbed state al, the transition state A1 results, yielding finally L-menthol. On the other hand, the transition state B1 and finally D-neomenthol result from the adsorbed state bl. The authors thank Prof. S. Imaizumi for helpful discussions. Partial financial support for this work by a grant for scientific research from the Japanese Ministry of Education is also acknowleged. We are most grateful to a referee for his helpful and scholarly comments. J. J. Villafranca and F. M. Raushel, Ado. Catal., 1979, 28, 323. K. Murakami, T. Harada and A. Tai, Bull. Chem. SOC. Jpn, 1980,53, 1356, 3367. T. Harada, Bull. Chem. SOC. Jpn, 1980, 53, 1919. Y. Saito and Y. Ogino, J. Catal., 1978, 55, 198. Y. Ogino, Catal. Rev. Sci. Eng., 1981, 23 (4), in press. Y. Saito, F. Miyashita and Y. Ogino, J. Catal., 1975, 36, 67. * K. J. Laidler, in Catalysis, ed. P. H. Emmett (Reinhold, New York, 1954), vol. I, p. 119. * A. Miyamoto and Y. Ogino, J. Catal., 1972, 27, 31 I. O W. Klyne and J. Buckingham, in Atlas of Stereochemistry: Absolute Configurations of Organic lo S . Glasstone, K. J. Laidler and H. Eyring, in The Theory of Rate Processes (McGraw-Hill, New York, Molecules (Chapman and Hall, London, 1974), pp. 30, 80 and 86. 1941), p. 350. (PAPER 1/599)

ISSN:0300-9599    

DOI:10.1039/F19827801079

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. Soc., Faraday Trans. 1, 1982, 78, 1087-1101 Proton-transfer Reactions in the Excited State of Phenanthrylamines by Nanosecond Spectroscopy and Fluorimetry BY KINZO TSUTSUMI, SHIZEN SEKIGUCHI AND HARUO SHIZUKA* Department of Chemistry, Gunma University, Kiryu, Gunma 376, Japan Received 22nd April, 198 1 Proton-transfer reactions in the excited singlet state of phenanthrylamines at 300 K have been studied by means of nanosecond time-resolved spectroscopy and fluorimetry. On the basis of dynamic analyses involving proton-induced fluorescence quenching, the proton dissociation (k,) and association (k,) rate constants and the acidity constants pK,* in the excited singlet state of phenanthrylamines were determined. The pK,* values are discussed theoretically by taking account of their electronic structures.An empirical equation to estimate approximately the correct pK,* values of aromatic amines using data from Stokes shifts and (pK,*),, values (determined by the Forster cycle) is proposed. Isotope effects on the rate constants and pK,* values have also been examined. Acid-base properties in the excited state of aromatic compounds are of interest in chemistry and also in biochemistry. They are closely related to the corresponding electronic structures. Since the original work of Forsterl and Weller2 a number of studies on the acidity constants pK,* in the excited state have been It is well-known that the pK,* value can be estimated by means of the Forster the fluorescence titration curve2, and the T, + T, absorbance titration curve? The Forster cycle involves the approximations that an acid-base equilibrium is established during the lifetime of the excited state and that entropy changes associated with proton dissociations both in the ground and excited states are the same.Recently the reliability of the Forster cycle has been questioned in some papers. Grabowski and Grabowskas reconsidered it with emphasis on the thermodynamic approximations and on the necessary experimental precautions. The titration curves also contain the assumptions that proton-transfer reactions in the excited state are very fast and that the acid-base equilibrium is established in the excited state. A laser study of the protonation equilibrium of triplet benzophenone has been reported by Rayner and Wyatt.lo They have shown that the pK,* value of the triplet benzophenone obtained by laser photolyses is almost equal to that estimated from the T, + T, absorbance titration curve, since the lifetime of the triplet benzophenone may be long enough to allow the acid-base equilibrium.However, for the excited singlet state it has been shown by Tsutsumi and Shizuka" that proton-induced fluorescence quenching competitive with proton-transfer reactions is present in the excited singlet state of naphthylamines and the simple acid-base equilibrium cannot be established in the S, state. Dynamic analyses involving proton-induced quenching are therefore needed in order to obtain the correct pK,* values in the S, state. Dynamic analyses by means of nanosecond time-resolved spectroscopy with fluorimetry were applied to P-aminopyrene,12 1 -aminoanthracene13 and naphthols.'* Recently, the Stuttgart group15 has supported our method of determining the pK,* values of naphthylamines.Similar studies on excited naphthols 10871088 PROTON-TRANSFER I N PHENANTHRYLAMINES have been reported by Harris and Selinger.16 The quenching mechanism induced by protons in polar media has been studied very recently by Tobita and Shizuka,17 and it is found that the fluorescence of aromatic compounds having an intramolecular charge-transfer structure is quenched effectively by protons and that the proton-induced quenching is caused by electrophilic protonation at one of the carbon atoms of the aromatic ring leading to proton exchange (or isotope exchange). As for phenanthrylamines, the acidity constants in the excited singlet and triplet states have been studied by means of the Forster cycle.l* The dynamic behaviour of the excited singlet state of phenanthrylamines in the presence of protons has been investigated in the present paper by means of nanosecond time-resolved spectroscopy with fluorimetry. The acid-base properties of the title compounds have been discussed from the viewpoint of n-electronic structures with the aid of a semi-empirical SCF-MO-CI method.EXPERIMENTAL The phenanthrylamines were synthesized except for 9-phenanthrylamine (9A) (Aldrich). Syntheses of 1 -phenanthrylamine (1A) and 4-phenanthrylamine (4A) were as follows. A mixture of p- 1 - and /3-2-naphthoylpropionic acids was prepared from naphthalene and succinic anhydride in nitrobenzene according to the method of Ha~0rth.l~ After each compound had been isolated by fractional crystallization, it was submitted to a Clemensen reduction to give 7-1- or y-2-naphthylbutyric acid, which was cyclized with concentrated H2S04 to produce 1- or 4-keto- 1,2,3,4-tetrahydrophenanthrene, respectively. These were then transformed to the oximes by the usual method, and then submitted to Beckmann rearrangement to give 1A [m.p.145-146 OC (146 "C20)] and 4A [m.p. 50-51 "C (55 "C20)], respectively. The overall yields were 1 and 2% for 1A and 4A, respectively. Syntheses of 2- (2A) and 3- (3A) phenanthrylamines were as follows. Phenanthrene was acetylated with acetyl chloride in the presence of aluminium chloride to give the mixture of 2- and 3-acetylphenanthrenes, which were separated by the method of Mosettig and Kamp.21 The acetylphenanthrenes were transformed to oximes with hydroxylamine hydrochloride, and were then submitted to Beckmann rearrangement22 to give 2A [m.p.85 "C (85 oC23)] and 3A [m.p. 86-87 "C (87.5°C23)]. The overall yields were 10 and 44% for 2A and 3A, respectively. 9A was purified by sublimation. H2S04 (97 %, Junsei), D20 (99.7 %, Merck) and D2S04 (96-98 %, isotopic purity 99 %, Merck) were used without further purification. Deionized water was distilled. Actual acid contents were determined by titration. The concentrations of the samples were 10-4-10-5 mol dm-3 in H20 (or D20) containing 10% acetonitrile in volume. In this study buffer solutions were not used since fluorescence quenching by inorganic anions might 25 All samples were thoroughly degassed by freeze-pumpthaw cycles on a high-vacuum line.Absorption and fluorescence spectra were measured at 300 & 1 K with Hitachi 139 and 124 spectrophotometers and a Hitachi MPF 2A fluorimeter, respectively. The fluorescence quantum yields at 280 nm excitation were measured by comparison with a quinine bisulphate 0.05 H2S04 solution (mF = 0.54).261 27 The fluorescence response functions were recorded at 300 & 1 K with a Hitachi nanosecond time-resolved spectrophotometer (1 1 ns pulse width). The kinetic analyses for the fluorescence response functions of A* and (A+H)* were carried out using an electronic computer (FACOM 230-25). METHOD OF CALCULATION Calculations were carried out on the basis of a semi-empirical SCF-MO-CI method according to the procedure described in a previous work.l* The C-C and C-N bond lengths were assumed to be 139 and 138 pm, respectively.All bond angles were assumed to be 120'. The computation was carried out with a HITAC 8800/8700 computer located at the Computer Centre of the University of Tokyo.K. TSUTSUMI, S. SEKIGUCHI AND H. SHIZUKA 1089 RESULTS AND DISCUSSION FLUORESCENCE TITRATION CURVES AND LIFETIMES IN THE EXCITED SINGLET STATE OF PHENANTHRYLAMINES Absorption and fluorescence spectra of the protonated and neutral phenanthryl- amines are shown in fig. 1. In the presence of protons ([H+] > mol dm-3), phenanthrylamines are protonated on the nitrogen atom in the ground state since the pK, values in the ground state of lA, 2A, 3A, 4A and 9A are 3.4, 4.1, 3.9, 3.18 and 3.5, respectively.18 The absorption spectra of the corresponding protonated amines A+H are similar to that of phenanthrene.There was no isotope effect on the shapes of the absorption and fluorescence spectra. On irradiation with 280 nm light, the neutral (A) and cation (A+H) fluorescence spectra were observed, de- pending upon proton concentrations. 3 A 15 20 25 30 35 15 20 25 30 35 h WY u .- E 15 20 25 30 35 15 20 25 30 35 15 20 25 30 35 wavenumber,’cm -* FIG. 1 .--(a) Absorption and (b) fluorescence spectra of neutral and protonated phenanthrylamines. lA, 2A, 3A, 4A and 9A denote I-, 2-, 3-, 4- and 9-phenanthrylamines, respectively. Logarithmic plots of the fluorescence quantum yields of neutral amines OA and cations (PAH as a function of [H+] are shown in fig.2. For example, the @A value for neutral 1A decreased considerably with increasing [H+]. In the range [H+] -c 1.5 mol dm-3, the OAH value increased slightly with increasing [H+]. Further- more, on over 1.5 mol dmP3 H,SO, being added, the QAH value increased appreciably with increasing [H+] to give a maximum value (PiH of 3 x lo-, at [H+] = 24.3 mol dm-3. Similar results for the fluorescence titration curves were also obtained in the case of FAR 1 36c TABLE 1 .-FLUORESCENCE QUANTUM YIELDS (mi,, a",), LIFETIMES (zi, z0AH), PROTON DISSOCIATION (k,), PROTONATION (k,) AND QUENCHING (kb) RATE CONSTANTS OF PHENANTHRYLAMINES AT 300 K sample" mi/ 10-2 zO,/ns khl k,/ a",/ 10-2 &/ns lo8 mol-l dm3 s-' kl/108 s-l lo8 mol-l dm3 s-I 1A H+ D+ 2A H+ D+ 3A H+ D+ 4A H+ D+ 9A H+ D+ 8.0, f 0.7 8.7, f 0.8 4.1 , +_ 0.4 4.2, f 0.4 4.5, f 0.5 4.5, _+ 0.4 4.8, f 0.5 4.6, f 0.2 4.7, f 0.3 3.86 f 0.3 8.5 f0.8 13.5 f0.9 13.5f0.5 14.4 f 0.5 10.7 k0.5 16.8f0.7 7.3k0.5 14.5+ 1 13.0 f 0.9 14.1 f0.5 2.9, f 0.3 4.4, f0.5 4-58 f 0.4 5.0, f0.3 6.2, f 0.3 6.9, f 0.4 4.4 f 0.4 3.1 f 0.3 3.3, +_ 0.4 3-38 f0.3 49f3 47., f 4 49., f 1 62., f 2 37f 1 52., f 1 49.1 f 5 37., f 4 35, k 1 38.,f 1 1.6 & 0.2 0.5, & 0.05 4.3, & 0.4 3.3,&0.3 0.93 & 0.08 0.44 & 0.04 2.4, k 0.3 2.1 & 0.2 1.9, & 0.2 1.6, k 0.2 8.1 f 0.9 6.2 f 0.9 9.2 f 0.6 7.7 f 0.5 7.5 f0.8 5.6 +_ 0.5 8.0 f 0.8 6.8 f 0.7 7.0 +_ 0.5 6.5k0.5 1.4k0.3 1.3 k0.3 1.1 kO.1 0.9 f 0.08 2.8 & 0.3 2.2 f 0.3 0.9 f 0.1 0.8 & 0.08 1.2 f 0.1 0.9 & 0.1 ~~ a For abbreviations see caption to fig.1. w P 0 cl 0 2 I clK. TSUTSUMI, S. SEKIGUCHI AND H. SHIZUKA 1091 - 3 - 2 - 1 0 1 2 - 3 - 2 - 1 0 1 2 - 1 - 2 - 3 - 3 - 2 - 1 0 1 2 log [H,O+l FIG. 2.-Logarithmic plots of the fluorescence quantum yields of (a) neutral amines (mA) and (b) cations (@AH) as a function of [H,O+]. 0, H,SO,+H,O; 0, D,SO,+D,O. other phenanthrylamines (fig. 2). The midpoints of the fluorescence titration curves of A* and (A+H)* differ, as can be seen in fig. 2. The discrepancy between the midpoints of the curves for @A and @AH indicates that for phenanthrylamines these curves do not correspond to simple two-component equilibria in the lowest excited singlet state. Therefore the pK,* values cannot be determined directly from the midpoints of the curves.These features are similar to those of naphthylamines," 1 -aminopyrene,12 1-amin~anthracenel~ and 1-naphth01.l~ The isotope effect on the fluorescence quantum yields of the phenanthrylamines is also shown in fig. 2. The fluorescence lifetime of neutral lA(zO, = 8.5 & 0.8 ns at ca. pH 7) decreased with decreasing pH. On the other hand, the apparent fluorescence lifetimes of the cation, rAH, were very short (rAH < 1 ns) in the range [H+] < 1.5 mol dm-3. At higher acid concentrations [H+] > 1.5 mol dm-3, rAH increased significantly with increasing [H+] to give a maximum value, z ; ~ , of 49+3 ns at [H+] = 24.3 mol dm-3. Similar results were also obtained for other phenanthrylamines.These data are listed in table 1. The isotope effect on the fluorescence lifetimes of phenanthrylamines is also shown in table 1. KINETIC ANALYSES The experimental results can be accounted for by the scheme shown in fig. 3,11 where k,, k, and kk denote the rate constants for proton dissociation, association and 36-21092 PROTON-TRANSFER I N PHENANTHRYLAMINES FIG. 3.-Reaction scheme of excited-state phenanthrylamines in the presence of protons. proton-induced quenching, respectively, kf and k; the radiative rate constants of (A+H)* and A*, respectively, and k, and kd the radiationless decay rate constants for (A+H)* and A*, respectively. From the steady-state approximation, the following equation can be obtained: Since the value of following equations should hold : is large and k, > k2 [H,O+] at [H,O+] c 0.5 mol drn-,, the and From the above relations, eqn (1) can be simplified to = 1 + k;zi[H,O+].@A A Stern-Volmer plot of @)"A/@, against [H,O+] gives a linear relationship (fig. 4), which agrees well with eqn (1'). The following experimental equations for phenanthrylamines were obtained : From the slopes of the linear plots and eqn (1') the rate constants for proton-induced quenching (kh) at 300 K were determined to be (1.6 f 0.2) x lo8 (lA), (4.36 f 0.4) x lo8 (2A), (9.3k0.8)~ lo7 (3A), (2.47k0.3) x lo8 (4A) and (1.92kO.2) x lo8 mol-l dm-3 s-l (9A), respectively. A deviation from linearity at higher proton concentrations ([H,O+] > 0.6 mol dm-,) for the phenanthrylamines was observed. This may be due to the variation in ionic strength at higher proton concentrations, and a more specific short-range interaction between A* and H+ seems to be dominant, as reported by Weller.28 In addition to this, the concentration of free water molecules, which act as proton acceptors of (A+H)*, should decrease at higher acid concentrations; the value of the pseudo-first-order rate constant kl(kl = k;[H20]) might also decrease underK. TSUTSUMI, S.SEKIGUCHI AND H. SHIZUKA 1093 . 0; 0 0.2 0.4 0.6 0 0.2 0.4 0.6 2 A 1 5 - 0 0.05 0.1 0 0.2 0.4 0.6 3 .O 2.0 I . o * r I I 0 0.2 0.4 0.6 [ H3 O+] mol dm-3 FIG. 4.-Stern-Volmer plots of against [H,O+]. 0, H,S04+H,0; e, D,SO, +D,O. such condition^.'^ Therefore, kinetic analyses involving proton-induced quenching should be carried out at moderate acid concentrations [H,O+] < 0.5 mol dm-3, where the value of k J k , is constant.mol dmA3 < [H,O+] d 5 x 10-1 mol drn-,), fluorescence from both A* and (A+H)* can be observed (fig. 2). Under a &function pulse excitation the fluorescence response functions FA(t) of A* and FAH(t) of (A+H)* are given by29-32 FA(t) = [k; kl/(A2 - A,)] (e-ll - e-4 t ) (7) and F’H(t) = [k,(A,-X)/(12-A,)J(e-.21t+ Ae-lZt) (8) At medium proton concentrations ( where A = ( X - A1)/(A2 - X). The decay parameters 1, and A, are A,,, = i ( X + Y T (( Y - a 2 + 4 k , k2[H,0+]}a) where x = (Z&)-l+kl and Y = (~;)-l+ (kd + k,) [H,O+]. (9)1094 PROTON-TRANSFER I N PHENANTHRYLAMINES z- 0 10 20 30 40 50 60 70 0 10 20 30 40 50 60 7 0 I0 20 30 40 50 60 70 0 10 20 30 40 50 60 70 0 10 20 30 40 50 60 70 tlns FIG. 5.-Decay curves of excited neutral amines in the presence of protons [(a) pulse, (b) ZA(t)].Solid line, computed; dashed line, observed. 1A: [H+] = 4.37 x lo-’ mol dm-3, 1, = 2.3 x lo8 s-l, I , = 1.33 x lo8 s-l; 2A: [H+] = 3.9 x lo-, mol dm-s, I , = 9.1 x lo7 s-,, 1, = 9.5 x lo8 s-l; 3A: [H+] = 7.59 x lob2 mol dm-3, A, = 9.83 x 107s-l, I , = 8.01 x loa s-,; 4A: [H+] = 1.31 x 10-l mol dm-3, I , = 1.67 x lo8 s-I, 1, = 8.41 x lo8 s-l; 9A: [H+] = 3.9 x lo-, mol dm-3, 1, = 8.4 x lo7 s-l, 1, = 7.34 x lo8 s-l. The output of the D, pulser is related to the undistorted fluorescence response function of FA(t) or FAH(t) by the convolution integral or where IL is the corresponding lamp function and IA(t) or IAH(t) the observed fluorescence response function of A* or (A+H)*, respectively. The observed fluores- cence response function IA(t) for the phenanthrylamines was analysed using eqn (7) and (10).A kinetic treatment of the fluorescence response function of (A+H)* was impossible since the intensity was very weak. Using the experimental values of 70,, 7 i H and kb the convolution method was applied, since exact values could be obtained in a convolution system containing a small number of parameters. Typical results for convolution in the H2S0, + H20 system are shown in fig. 5. Fig. 6 shows logarithmic plots of the decay parameters A1 and A2 as a function of [H,O+]. From the values of A1 and A2, both k, and k, values can be obtained. Similarly, kinetic analyses forK. TSUTSUMI, S.SEKIGUCHI AND H. SHIZUKA 1095 10- loo lo-* 10-l loo [H,O+] /mol dm-3 FIG. 6.-Decay constants A, and 1, as a function of [H,O+]. phenanthrylamines in the D2S04 + D20 system have been carried out. Isotope effects on proton-transfer reactions in the excited state may involve complex problems; those in the ground state have been discussed by Bell.33 A study of isotope and temperature effects on proton-transfer reactions in the excited state is in progress and will be reported in the future. PROTON-INDUCED FLUORESCENCE QUENCHING Appreciable proton-induced quenching of phenanthrylamines was observed, as described in the previous section. There was no quenching effect due to the counter-ion SO:- under the experimental conditions. It has been demonstrated for the quenching mechanism induced by protons that (1) intramolecular charge-transfer structure in the excited state is responsible for the proton-induced quenching and (2) the proton-induced quenching is caused by electrophilic protonation at one of the carbon atoms of the aromatic ring.17 In the excited state of phenanthrylamines charge migration from the amino group to the phenanthrene ring (ie.intramolecular charge-transfer) may take place in polar media just as it does in naphthylamines,ll 1 -naphthol,14 1 -methoxynaphthalene17 and 1 -aminopyrene.12 Electron migration in the excited state of aromatic amines is supported by the data on charge densities at the nitrogen atoms of the compounds calculated by a semi-empirical SCF-MO-CI method, as will be shown later (see table 2).However, the charge density at the carbon1096 PROTON-TRANSFER I N PHENANTHRYLAMINES TABLE 2.-ACIDITY CONSTANTS IN THE GROUND (p&) AND EXCITED SINGLET ( p g ) STATES AND 71-CHARGE DENSITIES ON THE NITROGEN ATOM OF PHENANTHRYLAMINES (PN) sample P C b 1A H+ 3.4 D+ 2A H+ 4.1 D+ 3A H+ 3.9 D+ 4A H+ 3.1, D+ 9A H+ 3.5 D+ - 0.76 f 0.08 - 0.68 f 0.05 - 0.92 & 0.08 - 0.93 f 0.08 - 0.43 f 0.04 - 0.40 f 0.04 - 0.95 f 0.09 -0.93 0.09 - 0.77 f 0.07 - 0.86 f 0.08 ~ ~~~~ -0.15f0.02 -5.6 1.8863 1.8370 0.15 * 0.02 -0.56k0.05 -2.6 1.8911 1.8352 -0.55f0.05 -0.01 fO.001 - 1.6 1.8903 1.8494 0.05 & 0.005 -0.45f0.05 -4.5 1.8863 1.8331 - 0.25 f 0.03 - 0.40 f 0.04 - 3.2 1.8864 1.9372 - 0.35 f 0.04 a Taken from ref. (18); determined by dynamic analyses; estimated from the fluorescence titration curve of A*.atom in the phenanthrene ring is not so great (< 1.085), since the migrating electron is widely distributed among the carbon atoms in the ring. Electrophilic protonation at one carbon atom in the aromatic ring (i.e. proton-induced quenching) is locally restricted.,' As a result, the values of kb for phenanthrylamines having such a small charge density on the carbon atom become small (ca. lo8 mol-l dm3 s-l) compared with those of 1-naphthylamine (ca. log mol-l dm3 s-l).ll The proton-induced quenching is competitive with protonation k , in the excited state of phenanthrylamines, as can be seen in table 1. It is therefore obvious that a simple acid-base equilibrium cannot occur in the S, state of phenanthrylamines. This is the reason why the fluorescence titration curves do not correspond to simple two-component equilibria in the S, state.For the excited state of protonated phenanthrylamines, (A+H)*, no proton-induced quenching was observed under experimental conditions. Protons are inactive to the cation (A+H)* owing to the electronic repulsion force between them. EXCITED-STATE p e VALUES OF PHENANTHRYLAMINES Dynamic analyses involving proton-induced quenching are needed in order to obtain the p e values of phenanthrylamines as well as those of naphthylamines.ll Using the values of k, and k, (table l), we can determine the correct p c values where the values of k l / k 2 are constant at moderate acid concentrations. The p c values are listed in table 2. Similarly, the p c values for the D,SO,+D,O system have been determined.The p c values in the excited singlet state of phenanthrylamines are more negative than those in the ground state. This well-known phenomenon is caused by electron migration from the amino group to the phenanthrene ring for the relaxed fluorescent state in polar media. In other words, the charge density or formal charge at the nitrogen atom of the excited singlet state of aromatic amines is closely related to basicity of the species, as will be discussed later. The pK,* values determined by the Forster cycle and by the fluorescence titration curve are also shown in table 2. The p e values obtained by the Forster cycle are very negative in comparison with those obtained by dynamic analyses. The fluorescence spectra of neutral phenanthrylamines in polar media shift considerably to the red, as shown in fig.1. The difference beween the p e values obtained by the Forster cycleK. TSUTSUMI, S. SEKIGUCHI AND H. SHIZUKA 1097 and those obtained by dynamic analyses arise mainly from the large Stokes shift of phenanthrylamines in aqueous media. The reliability of the Forster cycle has been discussed by Grabowski and Grabowskag on the basis of thermodynamic considera- tions. They have suggested that the Forster cycle may fail if the electronic levels of a given molecule invert during the lifetime of the excited state in a solvent-assisted relaxation process. Naphthylaminesll and 1 -naphthol14 are such cases, whereas no inversion of the electronic levels for phenanthrylamines takes place.However, the large Stokes shifts [(3.3-6.4) x lo3 cm-l] indicate a significant interaction between A* (lLa) and water molecules. In contrast, the p c values determined by the midpoint of the fluorescence titration curve of A* are positive in comparison with those obtained by the dynamic analyses. This discrepancy can be understood by taking account of the proton-induced fluorescence quenching, kb. ELECTRONIC STRUCTURES AND ACIDITY CONSTANTS OF PHENANTHRY LAMINES Electron migration from the amino group to the phenanthrene ring in the excited singlet state (lL,) occurs much more intensely than that in the ground state (lA). The charge densities or formal charges on the nitrogen atom of the phenanthrylamines have been estimated by means of a semi-empirical SCF-MO-CI method.The calculated results are shown in table 2. is applied to proton-transfer reactions in the excited state, a linear relation between the n-charge densities (P') or formal n-charges (el) and the excited-state p c values may be anticipated. Fig. 7 shows a plot of Qg If the chemical non-crossing 0.17 QN* QN -1.0 -0.8 -0.6 -0.4 -0.2 PK: 0.1 0.1 0.1 FIG. 7.-Plots of the acidity constants [pK,* (a) and pK, (b)] against the formal charges [QN and Q;], respectively.1098 PROTON-TRANSFER I N PHENANTHRYLAMINES (or QN) as a function of p c (or pKa), where the asterisk indicates the lowest excited singlet state. Calculations by the least-squares method give the following equations : pKa = - 143.4&+ 19.7 (r = 0.950) (12) and p c = -32.50Q;l;+4.50 ( r = 0.965) (13) where r denotes the correlation coefficient.state can be written as The Gibbs free-energy change (AG) for proton-transfer reactions in the ground (14) where Ka denotes the equilibrium constant. On the assumption that the difference in entropy changes between proton dissociation and association processes is small, the value of A H is given approximately by1* AG = - RT In Ka = (2.303 RT)pKa A H = I N - A + C (15) where I N is the valence-state ionization potential of the proper nitrogen atom, and A and C the electron affinity of hydronium ion and a constant value, respectively. Eqn (16) is derived from eqn (14) and (1 5) : IN +constant. (16) 1 2.303 RT - According to eqn (16), the theoretical pKa values should be proportional to the valence-state ionization potential of the nitrogen atom.In order to estimate the I N values (in ev), a parabolic relation between I N and the total valence-shell electron density (q) on the nitrogen atom is assumed:35 I N = a+ bq+cq2. (17) For the nitrogen, which is trigonal, a, b and c are equal to -89.402, 24.6265 and - 1.45 15, respectively. Eqn (1 7) can also be expressed as eqn (1 8) I N = -10.1622QN-2.1130 (18) where QN is the formal n-charge on the nitrogen atom and the term in QR is disregarded since the value of QK is very small. From a substitution of eqn (1 8) into eqn (1 6), eqn (19) is obtained : (10,1622QN + 2.1 130+ A - C) pK,theO= - 1 2.303 RT = - 1.71 x 102Q,+constant (19) where T = 300 K. Thus a linear relationship between pKkheo and QN holds, with a slope equal to - 1.71 x lo2.The slope of the line obtained experimentally is - 1.43 x lo2 for the ground state [eqn @)I. This value is almost the same as that derived by the theoretical method. For the excited singlet state, however, the slope of the straight line obtained experimentally is smaller than that of the theoretical relation for the ground state. This may be attributed to the relatively small dependence of in the S, state upon Q; compared with that of I N upon QN. H ~ y t i n k ~ ~ and also Waluk et have assumed a linear relationship between the logarithm of the protonationK. TSUTSUMI, S. SEKIGUCHI AND H. SHIZUKA 1099 I 1 1.83 1.84 1.85 1 hf ' 36 PN FIG. 8.-Plot of log k, against the n-charge density PN at the nitrogen atom. rate constant k, and the n-electron density P,.Fig. 8 shows a plot of log k, against P,, showing that the linear relation holds fairly well: log k, = 29.88PN -46.34 (r = 0.937). CORRELATION BETWEEN p e VALUES AND STOKES SHIFTS In the course of our studies on proton-transfer reactions in the excited singlet state of aromatic compounds, it has been shown that significant proton-induced quenching is present and that a simple acid-base equilibrium cannot occur in the S, state. Dynamic analyses involving proton-induced quenching are therefore needed in order to obtain the correct p e values. However, the treatment by dynamic analyses is complicated and troublesome in comparison with that by the Forster cycle. The large deviation of the values determined by the Forster cycle from the correct p c values is mainly due to the remarkable Stokes shift, as described above. The plot of A p e [ = p c - (pG)Fc] as a function of the Stokes shift AEst (in eV) is shown in fig.9, where the Stokes shifts were determined from the values of the peaks in the 10 8 6 4 2 n - 0 0.5 1.0 AEstleV FIG. 9.-Plot of A p e [ = p e - @K&c] values against the Stokes shifts (A&). Numbers denote 1, NNdimethyl-I-naphthylamine; 2, 1A; 3, 1-aminoanthracene; 4, 1-naphthylamine; 5, 4A; 6, 9A; 7, 2A; 8, 2-naphthylamine; 9, 1-aminopyrene; 10,3A. 0, H,SO,+H,O; 0, D,SO,+D,O.1100 PROTON-TRANSFER I N PHENANTHRYLAMINES absorption and fluorescence spectra. The following linear relation for aromatic amines is obtained by the least-squares method: A p e = 8.72AEst-2.64 (r = 0.885). (20) This plot also includes the A p e values of naphthylamines,ll NN-dimethyl- 1 - naphthylamine,ll 1 -aminoanthracene13 and 1 -aminopyrene,12 as have been previously reported.Fig. 9 shows that the values of A p e increase with increasing AEst: the AEst values increase with increasing solvation energy, caused by a re-orientation of the polar solvent molecules during the lifetime of the excited singlet state of the aromatic amines. As a result, an intramolecular charge-transfer state of the neutral excited amine is produced in polar media which is susceptible to electrophilic protonation (i.e. proton-induced quenching)l' to the aromatic ring. Using eqn (20') we can estimate approximately the correct p e values for the S, state of aromatic amines: p c = (pc)Fc + 8.7,AEst - 2.64.(20') CONCLUSION Proton-induced fluorescence quenching is involved in the excited singlet state of neutral phenanthrylamines, and a simple acid-base equilibrium cannot be established in the S, state. The relaxed fluorescent state, having an intramolecular charge-transfer structure in polar media, is susceptible to proton-induced quenching. In order to obtain the correct p e values, dynamic analyses involving proton-induced quenching are needed. Linear relations between the formal charges and acidity constants were obtained for both ground and excited singlet states independently. For aromatic amines the values of A p e [ = p e - ( P C ) ~ ~ ] are proportional to those of the Stokes shifts, i.e. Using this equation, the correct p e values can be estimated approximately from the data for ( p c & and AEst.p c = (pe)Fc + 8.72AESt-2.64. This work was supported in part by a Scientific Research Grant-in-Aid from the Ministry of Education of Japan. Th. Forster, Z . Elektrochem. Angew. Phys. Chem., 1950, 54, 42, 531. A. Weller, Ber. Bunsenges. Phys. Chem., 1952, 56, 662; 1956, 66, 1144. A. Weller, Progr. React. Kinet., 1961, 1, 189. E. Vander Donckt, Progr. React. Kinet., 1970, 5, 273. J. F. Ireland and P. A. H. Wyatt, Adv. Phys. Org. Chem., 1976, 12, 131. W. Klopffer, Adv. Photochem., 1977, 10, 31 1. (Wiley-Interscience, New York, 1966), p. 125. G. Jackson and G. Porter, Proc. R. Soc. London, Ser. A, 1961, 200, 13. Z. R. Grabowski and A. Grabowska, Z . Phys. Chem. (N.F.), 1976, 101, 197. ' E. L. Wehry and L. B.Rogers, Fluorescence and Phosphorescence Anulyses, ed. D. M. Hercules lo D. M. Rayner and P. A. H. Wyatt, J . Chem. Soc., Faraday Trans. 2, 1974, 70, 945. l1 K. Tsutsumi and H. Shizuka, Chem. Phys. Lett., 1977,52,485; Z . Yhys. Chem. (N.F.), 1978,111,129. H. Shizuka, K. Tsutsumi, H. Takeuchi and I. Tanaka, Chem. Phys. Lett., 1979,62,408; Chem. Phys., in press. l3 H. Shizuka and K. Tsutsumi, J . Photochem., 1978, 9, 334. l4 K. Tsutsumi and H. Shizuka, Z . Phys. Chem. (N.F.), 1980, 122, 129. l5 F. Hafner, J. Worner, U. Steiner and M. Hauser, Chem. Phys. Lett., 1980, 72, 139. l6 C. M. Hams and B. K. Selinger, J. Phys. Chem., 1980, 84, 891, 1366. l7 S. Tobita and H. Shizuka, Chem. Phys. Lett., 1980, 75, 140.K. TSUTSUMI, S. SEKIGUCHI AND H. SHIZUKA 1101 I" K. Tsutsumi, K. Aoki, H. Shizuka and T. Morita, Bull. Chem. SOC. Jpn, 1971, 44, 3245. Is R. D. Haworth, J. Chem. SOC., 1932, 1125. 2o W. Lagenbeck and K. Qeissenborn, Chem. Ber., 1939,72, 724. 21 E. Mosettig and J. van de Kamp, J. Am. Chem. SOC., 1930, 52, 3704. 22 W. E. Beckmann and C. H. Boatner, J. Am. Chem. SOC., 1936, 58, 2097. 23 A. Werner, Justus Liebigs Ann. Chem., 1902, 321, 312. 24 A. R. Watkins, J. Phys. Chem., 1974, 78, 2555. 25 H. Shizuka, T. Saito and T. Morita, Chem. Phys. Lett., 1978,56,519; H . Shizuka, M. Nakamura and 26 M. H. Melhuish, J. Phys. Chem., 1961, 65, 299. 27 J. N. Demas and G. A. Crosby, J. Phys. Chem., 1971,75, 991. 28 A. Weller, 2. Phys. Chem. (N.F.), 1955, 3, 238. 29 J. B. Birks, Photophysics of Aromatic Molecules (Wiley-Interscience, London, 1970); J. B. Birks, Organic Molecular Photophysics, ed. J . B. Birks (Wiley-Interscience, London, 1975), vol. 2, chap. 9. 30 B. Stevens, Adv. Photochem., 1971, 4, 161. 31 A. Weller, in The Exciplex, ed. M . Gorodon and W. R. Ware (Academic Press, New York, 1975). 32 W. R. Ware, D. Watt and J. D. Holmes, J. Am. Chem. SOC., 1974, %, 7853; M-H. Hui and 3 3 R. P. Bell, The Proton in Chemistry (Chapman and Hall, London, 1973); Chem. SOC. Rev., 1974, 3, 34 C. A. Coulson and H. C. Longuet-Higgm, Proc. R. SOC. London, Ser. A, 1947, 192, 16. 35 M. J. S. Dewar and T. Morita, J. Am. Chem. SOC., 1969, 91, 796. 38 G. J. Hoytink, Red. Trav. Chim. Pays-Bas, 1957, 11, 885. 37 J. Waluk, A. Grabowska and J. Lipinski, Chem. Phys. Lett., 1980, 70, 175. T. Morita, J. Phys. Chem., 1980, 84, 989. W. R. Ware, J. Am. Chem. SOC., 1976,98, 4718. 513; J. Chem. SOC., Faraday Trans. 2, 1980, 76, 954. (PAPER 1 /64 1)

ISSN:0300-9599    

DOI:10.1039/F19827801087

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. Soc., Faraday Trans. 1, 1982, 78, 1 103-1 1 15 Analysis of the Dependence on Temperature of Kinetic Solvent Isotope Effects Application to Kinetic Data for the Solvolysis of Organic Halides and Carboxylic Acid Anhydrides BY MICHAEL J. BLANDAMER,* JOHN BURGESS, NICHOLAS P. CLARE, PHILIP P. DUCE AND ROBERT P. GRAY Department of Chemistry, The University, Leicester LE 1 7RH AND Ross E. ROBERTSON Department of Chemistry, University of Calgary, Calgary, Alberta, Canada AND JOHN W. M. SCOTT Department of Chemistry, Memorial University, St John’s, Newfoundland, Camda Received 22nd April, 198 1 The implications are explored of both a single- and a two-stage mechanism for solvolytic reactions on the observed kinetic solvent isotope effect, ksie ( = k[D,O]/k[H,O]). The analysis identifies three types of behaviour for the dependence of ksie on temperature. In the first class, ksie changes regularly with temperature towards a limiting value.In a second class, ksie passes through either a maximum or minimum, the nature of the extremum depending on the difference between the apparent heat capacities of activation for reaction in the two solvents at the temperature of the extremum, Finally, and in exceptional cases, both a maximum and a minimum may be observed if the plots showing the dependence of apparent enthalpies of activation on temperature in the two solvents cross in two places. This analysis is used in an examination of the kinetic data for the hydrolysis of several organic carboxylic acid anhydrides where the case for a two-stage mechanism is well established.A similar mechanism for the solvolysis of t-butyl chloride, 2-chlorobutan- 1-01, 2,2-dibromopropane and 2-bromo-2-chloropropane clarifies the trends in previously reported activation parameters. The analysis reveals that no simple relationship exists between ksie and the apparent heat capacity of activation for solvolysis in water. For many years, Robertson1*2 advanced the hypothesis that the magnitude of the heat capacity of activation, AC;, and the kinetic solvent isotope effect, ksie (= k[D,O]/k[H,O]), for solvolytic reactions are linked to the extent of solvent re-organisation accompanying the activation process. Nevertheless, considerable controversy surrounds claims to the relative importance of the solvent isotope effects on the initial and transition ~ t a t e s .~ - ~ In these argument~l-~ a common assumption is that the solvolytic reaction involves a single stage with one kinetically important activation barrier. By way of contrast, Albery and Robinson3 suggested that, in the case of t-butyl chloride, the solvolytic reaction is a two-stage process where the first stage is reversible. In this case, the heat capacity of activation calculated from the dependence of rate constant on temperature and assuming a single-stage mechanism is an artefact of the analysis. Recently6 we have used the two-stage mechanism3 to account for the dependence of this heat-capacity term on solvent composition for the solvolysis of t-butyl chloride in a range of water + organic co-solvent mixtures.An extension of this analysis to the dependence of ksie on temperature forms the basis of this paper. 11031104 KINETIC SOLVENT ISOTOPE EFFECTS Examination of both single- and two-stage mechanisms leads to the identification of the possible patterns of behaviour in the dependence of ksie on temperature. We then consider the kinetic data'~ for the solvolysis of organic carboxylic anhydrides in water and D,O. These substrates undergo hydrolysis via a BAc2 mechanism' where the two-stage reaction scheme is formally identical to that proposed by Albery and Robinson3 for the solvolysis of t-butyl chloride. The data were originally analysed'? * to yield the activation parameters describing a single-stage reaction. It is therefore possible to compare and contrast the conclusions drawn from both analytical methods.A similar comparison can also be made between parameters describing the solvolysis of t-butyl chloride in waterg and D,O1° based on the two mechanisms. Indeed the values of ACZ for solvolysis of acetic anhydride and t-butyl chloride are similar., In addition, there is an indication that, for t-butyl chloride, ksie passes through a maximum near 292 K. This conclusion is based on a ratio of measured rate constants at roughly equal temperature^;^? lo e.g. at 285, 291 and 293 K, ksie = 0.705, 0.745 and 0.741, respectively. We also extend the treatment to include the data for 2,2-dibromopropane1°~ l1 and 2-bromo-2-chloropropane10~ l1 on the grounds that these solutes were previously classed as undergoing solvolysis by an SN 1 mechanism, a conclusion prompted by the large negative values for ACg.We also include one system, 4-chlorobutan- 1 -o1,l29 l3 where there is a possibility of neighbouring group participation.12 The results of the analysis lead to the conclusion that the value of ksie at some arbitrary temperature is not tremendously informative. However, an analysis of the data using a two-stage mechanism pinpoints those features which may lead to extrema in kSie. ANALYSIS The experimental quantity is the first-order rate constant k at temperature T and ambient pressure. The mechanism for the reaction is described in terms of two models and related reaction schemes. We then consider the ratio k/ko, where ko is the rate constant for reaction in water and k the rate constant for reaction in D,O, i.e.ksie = k/ko. MODEL 1 The rate constant k characterises a single-stage reaction. Previously' the dependence (1) It follows1 that ACZ is calculated from the estimate Ci,, and AH# at temperature T from 6, and Ci,. Implicit in eqn (1) is the assumption that AC,Z is independent of temperature. In an application of eqn (1) to a set of data it is necessary to test the statistical significance of including the third term; if it is not significant, In k is simply expressed as a linear function of T-l to yield AH# as the slope. The dependence of In (kSie) on temperature is calculated from the two sets of parameters, ai and a: where i = 1, 2 and 3. According to eqn (l), In (ksie) has an extremum at one temperature: (2) of k on temperature was fitted to eqn (1) using a linear least-squares procedure: In k = a, +a, T-l+ a3 In T.T(kSi, = extremum) = (a, - a9/(a3 -a:). The solvent isotope effect on the derived activation hparameters is represented by the medium operator,14 6,; e.g. eqn (3) for the heat capacity terms Thus at the temperature given by eqn (2), 6 , A W = 0.BLANDAMER et al. 1105 MODEL 11 The two-stage mechanism for the reaction is written3 as shown in eqn (4) k, k3 reactant (intermediate) + product. k2 If a = k,/k3, the observed rate constant is related to k , and a by eqn ( 5 ) (4) k = kJ(1 -a). ( 5 ) It is assumed3? 6, l5 that both In k, and In a are linear functions of T-l. Previouslys we showed how the dependence of k on Tcan be fitted to eqn (6) using a non-linear least-squares technique : k = x, exp (x,/ T ) / [ 1 .O + x3 exp (x4/ T)]. Therefore AH? is calculated from x, and AAHZ (= AH$ -AH:), from x4.(Both AH? and AAH# are sufficiently large that we may overlook the small curvature in plots of In k , and In a against T-l required by classical transition-state theory in the event that AH? and AAHZ are independent of temperature.) Thus a positive value for AHf means that In k , increases with increase in temperature whereas a positive AAHZ means that a decreases. The kinetic solvent isotope effect is given by eqn (7) Thus ksie is dependent on the ratio of (1 +ao) to (1 +a) rather than ao to a. This is a source of complexity. By fittings the dependence of k(D20) on T and then k(H20) on T to eqn (6), it is possible to calculate (k,/ky) and (a/ao) together with the solvent isotope effects on AH? and AAH#.COMPARISON OF MODELS 1 A N D 11 Here we contrast the two models for the solvolysis reaction and examine those conditions which lead to a maximum and/or minimum in the ksie. We recall that such an extremum is observed for acetic anhydride.'* The starting point is the equation3. 6* l5 relating the enthalpy of activation AH# defined by model I and the two enthalpy terms, ACg1 and AAH#, introduced in model 11: (8) AH# = AH? +[a/(l +a)]AAH#. Eqn (8) describes a sigmoidal c ~ r v e ~ ~ ~ with a point of inflexion near the temperature where a = 1.0. In the context of model 11, the shape of this curve plays an important r61e in determining the dependence of ksie on temperature. Complexities emerge because a non-zero value for AAH# requires that a is dependent on temperature. Therefore the heat capacity of activation according to model I is given,6* l5 in terms Consequently AC; is negative irrespective of the sign of A A H Z , the dependence of AC; on temperature forming an inverted bell-shaped plotsv l5 with a minimum close to the temperature where a = 1.0.According to model I, the condition for In (kSie) to show an extremum is given by1106 KINETIC SOLVENT ISOTOPE EFFECTS The extremum is observed at the temperature where AH# = AHo#. Clearly if these quantities differ and are independent of temperature, no extremum will be observed. However, in terms of model 11, a more complicated pattern emerges. The analogue of eqn (10) is given by eqn (1 1) If the term in braces on the right-hand side of eqn (1 1) is independent of temperature, then as the temperature increases, In (k/ko) will either increase to an asymptotic limit if the term is positive or decrease to a lower limit if this term is negative.If both a and ao 6 1 , the overall trend is determined by the difference between AH? and AH:#. If both cc and ao >> 1, the trend in ksie is determined by the relative signs and magnitudes of four enthalpy terms, AH?, AH:#, AAH# and AAH*#. However, further complications emerge when either a or ao or both are, in terms of magnitude, near unity. It is informative to differentiate eqn (1 1) with respect to temperature to obtain eqn (12): d2 In dT2 (ksie) = -L{[AHf RT3 + ( & ) A A H # ] - [ A H : # + ( ~ ) A A H o # ] ] 1 +ao +- RT2 (l+a0)2 RT2 1 Eqn (1 1) shows that an extremum in ksie will be observed at the temperature where the sigmoidal curves showing the dependence on temperature of AH# [cf.eqn (S)] intersect. (There is no a priori requirement that they do so; no intersection occurs if, at all temperatures, AH# > or < AHo#.) At the temperature of intersection, the second differential, eqn (12), can be written as shown in eqn (13) where we have incorporated eqn (10) : If, therefore, ACpZ < ACo# at the intersection of the enthalpy curves, ksie is a maximum; if AC,Z > ACF, ksie is a minimum. A point of inflexion occurs where the second differential, eqn (1 2), is zero. No simple relationship between activation parameters emerges under this condition, although such a feature will be favoured near the temperature at which the plots showing the dependences of ACZ and ACif on temperature [eqn (9)] intersect [cf.eqn (13)]. The essential features of the analysis are summarised diagrammatically in fig. 1. Here the dotted line represents the reference system, the enthalpy plots crossing at A and B and the heat-capacity plots crossing at C . It follows from the values of ACZ and AC;# at temperature A, that here ksie is a minimum whereas at temperature B, ksie is a minimum. A point of inflexion in the dependence of In (Itsie) on temperature will occur near temperature C. These features will only be observed if the experimental temperature range includes A, B and C. In certain systems only one extremum may be observed and in others none at all. In the latter case, this might arise if the experimental range does not include the temperatures A, B or C.In certain systems, a point of inflexion may occur in the absence of an extremum but ,near the temperature where the two enthalpy curves come close together without actually crossing [cf. eqn (1 3 1 -BLANDAMER et al. 1107 B temperature 1 temperature - FIG. 1.-Schematic diagram showing the dependence of enthalpies and heat capacities of activation on temperature for solvolysis in water (full line) and deuterium oxide (dotted line). CALCULATIONS Numerical analysis of the data was carried out using computer programs (FORTRAN) written for the CDC Cyber computer at the University of Leicester. The linear least-squares fitting of the kinetic data to eqn (1) used a statistical packagels (GLIM).The non-linear least-squares fitting of the kinetic data to eqn (6) used the technique previously described.s RESULTS In this section we comment in some detail on the analysis of the data for the solvolysis of acetic anhydride’ and t-butyl chloride.@? lo However, we draw attention to the results of the analysis for other systems. At better than a 95% confidence level (Student’s ?-test), it is significant to include the third term in eqn (1) with reference to the kinetic data for solvolysis of acetic anhydride in both water and D20. The corresponding values of AC$ (i.e. model I) are negative, being strikingly more negative in water than in D,O (table 1). The maximum in In ksie is predicted [eqn (2)] to occur at 288.7 K, which is within the experimental range [cf.table 2 of ref. @)I. The solvent isotope effect on AH# (e.g. at 290 K, d,AH# = 1.16 kJ mol-l) is small but the isotope effect on AC$ is dramatic; S,ACg = 127 J mol-l K-l (table 1). A similar endothermic change occurs for propionic anhydride together with an appreciable change in AC$ but, in this case, the latter is in the opposite direction; SmAC: = -200 J mol-l K-l. A similar comparison is unfortunately not possible for the remaining three anhydrides because, as shown in table 1, these include systems where it was not possible to estimate u3 and at to the required confidence level. It is, however, noteworthy that for both phthalic and benzoic anhydrides, AC$ for solvolysis in D20 is large and negative. The result of the analysis for t-butyl chloride (table 1) shows a similar endothermic1108 KINETIC SOLVENT ISOTOPE EFFECTS change in AH# on going from water to D20.A similar pattern emerges for the remaining three systems, as indeed is the change in ACZ to a more negative value, e.g. for t-butyl chloride, S,AC,Z = - 146 J mol-1 K-l. The data predict that In ksie is a maximum when T = 303.4 K, which is just above the experimental temperature range. Thus when T < 303.4 K, S,AHZ > 0 and when T > 303.4 K, &,AHf < 0. When T = 303.4 K, ksie = 0.7514. Note that for 2-bromo-2-chloropropane SmAC,f is essentially zero, bearing in mind standard errors on each ACZ term. TABLE 1 .-DERIVED ACTIVATION PARA METERS^ (MODEL I) FOR SOLVOLYSIS OF ACID ANHYDRIDES AND ORGANIC HALIDES IN WATER AND DEUTERIUM OXIDE solute A H z (298 K)/ -ACz/ solvent ref.kJ mol-l J mol-l K-l acetic anhydride propionic anhydride succinic anhydride phthalic anhydride benzoic anhydride t-butyl chloride 2,2-dibromopropane 2-bromo-2-chloropropane 4-chlorobutan- 1-01 7 7 7 7 8 8 8 8 8 8 9 0 1 0 1 0 2 3 40.24 (0.18) 41.40 (0.15) 39.68 (0.07) 40.29 (0.23) 55.12 (0.45) 55.24 (0.12) 59.35 (0.57) 53.65 (0.77) 46.45 (0.32) 47.43 (0.46) 94.37 (0.28) 95.13 (1.16) 112.73 (0.13) 1 16.53 (0.18) 108.54 (0.02) 11 1.37 (0.06) 101.12 (0.33) 105.78 (0.28) 311 (14) 184 (29) 144(11) 344 (44) b b b 207 (62) 219 (27) 348 (19) 348 (20) 453 (16) 394 (8) 419 (33) 172 (8) 249 (7) b 494 (94) a Shown in brackets are the standard errors on derived quantities; third term in eqn (1) not significant at the 95% confidence level (Student's t-test).The results in table 1 follow from application of eqn (1) to the data. In the five cases where both SmAC,Z and S,AHf could be estimated, the analysis is satisfactory as expressed by the closeness of the fit between observed and calculated rate constants. Nevertheless some features of the analysis were disturbing. In those cases where it was statistically acceptable to use three terms in eqn (1) the correlation coefficients between the estimates of ai quantities (calculated from the normalised variance-covariance matrix) were in all instances close to unity, i.e. 1 .OO. We have commented elsewhere" on the unsatisfactory features of eqn (1). Nonetheless the analysis does identify important features of the data when analysed on the basis of model I.Turning to the analysis based on model 11, the data for the solvolysis of acetic anhydride in both water and D20 can be satisfactorily fitted to eqn (6). For the 36 individual rate constants describing reaction in water, agreement between calculated and observed values of k was within +0.78%, a plot of the residuals against temperature showing a satisfactory random scatter. Consequently it would not be justified to extend eqn (6), thereby leading to calculation of the dependence of AH:TABLE 2.-KINETIC SOLVENT ISOTOPE EFFECTS; ANALYSIS IN 'TERMS OF MODEL 11 acetic anhydride t-butyl chloride 2,2-dibromopropane 2-bromo-2-chloropropane 4-chlorobutan- 1-01 H 2 0 D20 H 2 0 D20 H2O D2O H2O D2O H2O D2O n data points T range/K k, (290 K)/10-2 s-l AHt/kJ mol-1 a (290 K) -AAH#/kJ mol-1 T(a = l.O)/K -AC,f (min)/J mol-' K-l T(ACpZ = min)/K -AC,Z (290 K/J mol-l K-I ksie (a/ao; 290 K) fit, A/% ksie (k,/k;; 290 K) 36 42 20 37 275-298 278-3 13 274-293 277-294 0.78 2.8 0.25 3.4 1.705 1.004 1.126 1.225 86.95 87.89 106.38 121.54 9.535 17.097 0.1525 0.691 46.46 45.66 51.98 51.43 259.6 252.2 317.7 295.1 961 994 813 914 255 250 315 290 265 156 444 914 0.589 1.088 1.79 4.53 1 43 35 13 48 293-3 18 295-323 283-308 2863 13 2.4 2.7 2.5 3.7 1.287 x loT3 0.903 x lop3 5.162 x 3.556 x 1 15.33 118.44 114.97 1 18.29 62.47 61.77 60.65 60.05 349.1 346.5 335.0 333.58 970 962 988 986 340 345 330 330 68 81 167 185 0.012 0.015 0.034 0.039 0.702 0.689 1.25 1.14 42 14 323-358 318-358 1.4 3.4 5.503 x 3.989 x lop3 98.26 100.61 2.92 x 5.087 x 70.14 68.89 402.66 394.7 920 922 400 390 2 3.5 0.725 1.7421110 KINETIC SOLVENT ISOTOPE EFFECTS and AAH# on temperature. A similar state of affairs exists for the data describing the solvolysis of acetic anhydride in D,O, agreement between observed and calculated rate constants being within f 1.7%.From both sets of derived parameter [eqn (6)], the related kinetic parameters have been calculated (table 2) together with the dependences of AH#, AHo#, ACZ and AC;# on temperature [eqn (8) and (911; fig. 2. The minimum in AC,Z for solvolysis in D,O is at a slightly lower temperature 30- 3 50 250 300 TI K FIG. 2.-Solvolysis of acetic anhydride in water (full line) and D,O (dotted line). than for solvolysis in water. The minimum is also marginally more intense.The curves showing the dependence of AH# and AHof on temperature cross when T = 291 K. At this temperature (fig. 2), AC;# < AC$, a feature reflected in the corresponding minimum in In kSie. In the high-temperature limit, AH# c AHo# and so In ksie increases with increase in temperature. For comparison we have selected a common temperature, 290 K, and calculated the isotope effects on k,, a, AH? and AAH#, table 2. In this system, the separate kinetic and activation parameters for the two- stage process in each solvent combine to yield a minimum in In ksie within the measured temperature range [fig. 2(c)]. The analysis of the data for the remaining anhydrides was less satisfactory. The 40 data points for the solvolysis of propionic anhydride in water were satisfactorily fitted to eqn (6) but such was not the case for the 24 data points for solvolysis in D20.Despite repeated attempts to locate a minimum in the least-squares analysis, agreement between observed and calculated rate constants was in some cases no better than + 6%. This failure was disappointing because In ksie passes through a maximum, thereby contrasting propionic and acetic anhydrides. The kinetic data for phthalic and benzoic anhydrides in D,O were satisfactorily fitted to eqn (6), but such was not the case for the data describing solvolysis in water and for the kinetic data for succinic anhydride in both water and D20. The results of the analysis using eqn (6) for t-butyl chlorideQ* lo are summarised in table 2. For solvolysis in water agreement between observed and calculated rateBLANDAMER et ai.1111 -200 -400 - k I L E -600 - --- iF. Q - 800 1 2 0 k 2 60 - 300 T/K 1.C 0 n 0 -Y 2 -1-0 v c - -2.0 2 - 3 50 TI K FIG. 3.-Solvolysis of t-butyl chloride in water (full line) and D,O (dotted line). -1000 300 350 TIK I 3 50 1 TIK 260 300 3b0 TI K FIG. 4.-Solvolysis of 2-bromo-2-chloropropane in water (full line) and D,O (dotted line).1112 0.2 0 - -0.2 5 0 - --. -Y - G - - 0.4 -0.6 KINETIC SOLVENT ISOTOPE EFFECTS I 1 1 1 280 300 320 340 31 TIK ( C ) 0 1 I I 2 80 300 320 340 3f TIK FIG. 5.-Dependence of (A) In (k,/k:) and (B) In (ala") on temperature for (a) acetic anhydride, (b) t-butyl chloride, (c) 2-bromo-2chloropropane, ( d ) 2,2-dibromopropane and (e) Cchlorobutan- 1-01.BLANDAMER et al. 1113 constants was within +0.25% for 20 data points while for solvolysis in D20 the agreement was within & 3.0% for 37 data points.The calculated dependences of A H #, AHo#, AC:, AC$# and In ksie on temperature are shown in fig. 3. For this system there is a dramatic solvent effect on the temperature corresponding to the minimum in AC,Z [fig. 3 (a)]. The two curves cross at 307 K, whereas the curves for the enthalpy terms cross at 292 and 321 K, features producing a maximum and a minimum, respectively, for In ksie; fig. 3(c). However, the overall trend is for In ksie to increase with increase in temperature. A similar trend is observed for 2-bromo-2-chloropropane (fig. 4). However, for this system the corresponding enthalpy plots do not intersect so there is no extremum in In ksie although there is a point of inflexion near the temperature where the heat-capacity curves intersect (fig.4). This pattern is also followed by 2,2-dibromopropane. The corresponding plots for 4-chlorobutan- 1-01 resemble those for t-butyl chloride (fig. 3) except that the enthalpy plots intersect at 357 (where In ksie is a maximum) and the heat-capacity plots intersect at 397 K. The most important set of results are summarised in fig. 5, where we show the calculated dependence on temperature of the kinetic isotope effects for k , and a, i.e. In (k,/ky) and In (a/ao), over the experimental temperature ranges. DISCUSSION The established two-stage mechanism for the solvolysis of acetic anhydride provides a sound basis for analysis of the kinetic solvent isotope effect using model 11.Further, it is possible to identify the relationships between the thermodynamic activation parameters calculated from models I and I1 for the same set of kinetic data. For example, the increase in AC,Z for acetic anhydride on going from water to D20 (table 1) is according to model I1 misleading. It is not, for example, indicative of a dramatic difference in solute-solvent interactions in water and D20. According to model 11, there is only a slight change in this now apparent AC,f at the minimum (fig. 2). However, when this is combined with a modest shift in the temperature at this minimum (table 2), there is a more marked effect on some averaged AC,Z quantity calculated using eqn (1). If this interpretation of the trends in previously reported AC$ parameters is correct, then the information yielded by changes in A C g on going from water to D20 is not straightforward because AC,Z is a function of AAH# and a [eqn (9)].Indeed no simple relationship exists, according to model 11, between AC,Z for solvolysis in water and either the solvent isotope effect at one temperature [cf. eqn (7)] or the dependence of ksie on temperature [cf. eqn (1 l)]. Thus in terms of model 11, In ksie and its temperature dependence is a complex quantity reflecting the solvent isotope effect and temperature dependence of k , and a; fig. 5. Consequently the extremum in In ksie for acetic anhydride has no particular significance, except in so far as the experimental temperature range includes the temperatures where a and ao are most sensitive to temperature.Problems emerge as shown for other systems (cf. table 2), where following a fitting of the data for solvolysis in water and D20 separately to eqn (6), the analysis predicts extrema and points of inflexion in ksie outside the experimental range. However, we would argue that from inspection of the curves given, for example, in fig. 2-4 drawn over an extended temperature range, it is possible to identify those features which contribute to the dependence of In ksie over the measured temperature range. For example, in the case of acetic anhydride in either water or D20, AAH# is negative and so a and ao increase with increase in temperature. Therefore, the high-temperature limit of A H # (model I) is AH? - IAAH 21.On going from water to D20, the high-temperature limit of A H # is displaced vertically since d,AH,' > d,AAHZ > 0 (table 2). If all other quantities were unaffected by the1114 KINETIC SOLVENT ISOTOPE EFFECTS change in solvent of if T(a = 1) has also increased, no extremum in ksie would be observed. However, the temperature T(a = 1) decreases, resulting in an intersection of the enthalpy plots (fig. 2). The rather poor analytical results for propionic anhydride indicate that similar small changes in AH? and AAH # can produce a maximum in Inksie. For the other anhydrides (table 1) the kinetic data cover too small a temperature range to allow calculation of the trends in ksie based on model 11. If, however, for benzoic and phthalic anhydride in water, AH# is independent of temperature, the corresponding dependence (model 11) for AH # in D20 intersects this line producing extrema as previously reported.8 In a similar fashion to that for the acetic anhydrides, it is apparent that the analysis in terms of model I1 resolves many of the problemslO in interpreting the solvent isotope effects for the solvolysis of organic halides using model I.Thus the dramatic effect on ACZ for t-butyl chloride (table 1) is a combination of a decrease in AC$ at the minimum combined with an increase in the temperature at the minimum. In contrast, the small change in AC$ (model I) for 2-bromo-2-chloropropane is a consequence of a less dramatic solvent effect on AC$ (min) and the temperature corresponding to this minimum.Indeed, in contradiction to the claimed generalisation,1° AH # is not necessarily larger for solvolysis in D20 than in H20. Even if eqn (1) is retained, it was not previously recognized that the derived parameters, ai in eqn (l), predict that In ksie passes through a maximum. Certainly, in terms of model I1 there are indications of a more clear-cut set of generalisations. As shown in fig. 5, k/ko for five substrates increases with increase in temperature whereas a/ao decreases; the latter more gradually. Significantly the limiting high-temperature limit for both ratios is not unity.3 In all systems the temperature where a = 1.0 increases on going from water to D20, but there is not dramatic change in A A H f . The latter shows that the differences in the two enthalpies of activation governing the fate of the intermediate [eqn (4)] is insensitive to the change in solvent.It is, however, not unexpected that S,AH,Z > 0. Thus if the formation of the transition state involves charge-separation, we would anticipate that the partial molar enthalpy of the transition state increases on going from water to D,O because enthalpies of transfer of salts are, with the expection of large alkylammonium salts, generally positive.18 This trend combined with a decrease in the partial molar enthalpy of a hydrophobic initial state19 produces a positive value 6 , A H t . The fact that the derived parameters for 4-chlorobutan-1-01 fall into the same general pattern (table 2) is in agreement with the previous conclusion12 that the extent of neighbouring-group participation in the activation processes is small.Further, the overall self-consistency between the solvent isotope effects on the parameters based on model I1 lends support to the proposal made by Albery and Robinson3 concerning the operation of a two-stage mechanism. Nevertheless, we must conclude that the ksie for a given reaction at one temperature is not a useful guide to the mechanism of reaction. We thank the S.R.C. for a maintenance grant to P.P.D. We thank the Royal Society for a travel grant to M.J.B. 3 4 1 2 R. E. Robertson, Prog. Phys. Org. Chem., 1967, 4, 213. P. M. Laughton and R. E. Robertson, Solute-Soluent Interactions, ed. J. F. Coetzee and C . D. Ritchie (M. Dekker, New York, 1969), chap. 7. W. J. Albery and B. H. Robinson, Trans. Faraday SOC., 1969, 65, 980; 1623. C. G. Swain and E. R. Thornton, J. Am. Chem. SOC., 1962, 84, 822. P. M. Laughton and R. E. Robertson, Can. J. Chem., 1965, 43, 154. M. J. Blandamer, J. Burgess, P. P. Duce, R. E. Robertson and J. M. W. Scott, J. Chem. SOC., Faraday Trans. I , 1981, 77, 1999.BLANDAMER et at. 1115 R. E. Robertson, B. Rossall and W. A. Redmond, Can. J. Chem., 1971, 49, 3665. B. Rossall and R. E. Robertson, Can. J. Chem., 1975, 53, 869. E. A. Moelwyn-Hughes, R. E. Robertson and S. E. Sugamori, J. Chem. Soc., 1965, 1965. A. Queen and R. E. Robertson, J. Am. Chem. Soc., 1966,88, 1363. lo L. Treindl, R. E. Robertson and S. E. Sugamori, Can. J. Chem., 1969,47, 3397. l2 M. J. Blandamer, H. S. Golinkin and R. E. Robertson, J. Am. Chem. Soc., 1969, 91, 2678. l3 H. S. Golinkin and R. E. Robertson, unpublished data. l4 J. E. Leffler and E. Grunwald, Rates and Equilibria of Organic Reactions (Wiley, New York, 1963). Is J. M. W. Scott and R. E. Robertson, Can. J. Chem., 1972,50, 167. l7 M. J. Blandamer, J. M. W. Scott and R. E. Robertson, Can. J , Chem., 1980, 58, 772. l8 H. L. Friedman and C. V. Krishnan, in Water: A Comprehensive Treatise, ed. F. Franks (Plenum l9 G. C. Kresheck, H. Schneider and H. A. Scherago, J. Phys. Chem., 1965,69, 3132. GLZM (Royal Statistical Society, London, 1977). Press, New York, 1973), chap. 1. (PAPER 1/648)

ISSN:0300-9599    

DOI:10.1039/F19827801103

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1, 1982, 78, 11 17-1 126 Seeded Emulsion Polymerizations of Styrene The Fate of Exited Free Radicals BY BARRY C. Y. WHANG, DONALD H. NAPPER,* MATTHEW J. BALLARD AND ROBERT G. GILBERT Departments of Physical and Theoretical Chemistry, University of Sydney, New South Wales 2006, Australia AND GOTTFRIED LICHTI Australian Atomic Energy Research Establishment, Sutherland, New South Wales 2232, Australia Received 7th May, 1981 The kinetics of the seeded emulsion polymerization of styrene at different particle number concentrations and different initiator concentrations have been studied. The results were analysed using a theoretical treatment that incorporates the possibilities of both the re-entry ofexited free radicals into the latex particles and the cross-termination in the aqueous phase of exited free radicals by free radicals generated through initiator decomposition.The results exclude the possibility of complete re-entry of exited free radicals into the latex particles for the initiator concentrations studied. They strongly support the occurrence of significant cross-termination in the aqueous phase of the exited free radicals with free radicals generated by initiator decomposition. This is in accordance with the known rapidity of cross-termination reactions compared with the corresponding self-termination reactions. It was also shown that the rate of entry of free radicals into each latex particle increases with decreasing particle number at constant initiator concentration. The radical capture efficiency was, however, relatively insensitive to the particle number concentration. The overall polymerization rate was found to be a complex function of the number concentration of latex particles; this is because this overall rate is itself a complicated function of the rate coefficients for entry, exit, etc.each of which may be individually a simple function of number concentration. Such behaviour reflects significant deviations from Smith-Ewart case 2 kinetics that occur in these sysems, rather than deviations from the general Smith-Ewart kinetic scheme. Previous studies'. of seeded emulsion polymerizations of styrene showed that free- radical exit (desorption) could play an important role in determining the kinetics of polymerization. In those studies, however, for all except very low initiator concen- trations, the exit rate was only a small fraction of the rate of entry of free radicals into the particles.Consequently, the results cast little light on the importance of the re-entry of exited free radicals into the latex particle^.^-^ The role of radical re-entry can be explored by increasing the relative importance of the exit rate and/or varying the number concentration of the seed latex particles. The former can be accomplished in two different ways: first, by the use of relatively small seed latex particles, since it was shown previously1 that the exit rate coefficient ( k ) is inversely related to the latex particle size; secondly, by the reduction of the initiator concentration so that the entry rate coefficient (p) is decreased.l In what follows, the results of these types of experiments are presented. The experiments highlight the importance of cross-termination of exited free radicals in the aqueous phase.They also confirm that the rate of emulsion polymerization is not always 11171118 SEEDED EMULSION POLYMERIZATION directly proportional to the number of latex particle^.^ An explanation for this type of behaviour in terms of the precepts of the Smith-Ewart theorys is proposed. EXPERIMENTAL The preparation and characterization of the monodisperse polystyrene seed latex have been described in detail previously.' Its mean particle radius was 38.3 nm, comparable to that of latex R 12/ 150 reported previously. mote that the latex designations R 12/ 150 and R 12/ 17 1 were inadvertently interchanged in table 1 of ref.(l).] The dilatometric method for measuring the rate of polymerization at 50 O C has also been detailed elsewhere.' Results are presented only for those systems that exhibited no particle nucleation as determined by electron microscopy. We note parenthetically that the induction period was found to increase significantly with increasing particle number concentration. Values for the contraction factor, the propagation rate constant and the equilibrium concentration of monomer in the latex particles were determined previously.' THEORETICAL TREATMENT OF EXPERIMENTAL DATA SLOPE A N D INTERCEPT METHOD FOR MODELS INCORPORATING RADICAL RE-ENTR Y AND C ROSS-TER MI NATION We have previously1 presented a rigorous theoretical analysis of the approach to a steady state in an emulsion polymerization for interval I1 that is accurate provided that the average number of free radicals per particle (A) does not exceed 0.5.This treatment (which did not take re-entry specifically into account) must of course only be used in such a limit and cannot be extrapolated to values of A greater than 0.5.' This analysis permits values to be obtained forp and k from the slope of the steady-state rate line and its intercept. This analysis did not, as already stated, include the possibility of re-entry of radicals that had exited from the latex particles. In what follows, we incorporate the possibilities of re-entry and cross-termination into the model so that it is possible in this case to transform measurements of the steady-state slope (a) and intercept (b) into values of p and k .Note that the intercept can be either negative (as in chemically initiated seeded systems1) or positive (as in relaxation studies2). We have shown previously1 that this treatment, which we term the 'zero-one' approximation (since only particles containing 0 or 1 free radicals are considered), is valid if the termination rate coefficient c % k,p so that the average number of free radicals per particle is not large ( A < 0.5). The kinetic equations incorporating the possibility of re-entry and cross-termination are -- dNo - - (PA 4- akA)No 4- (PA 4- akii+ k)Nl dt where N , is the relative population of particles containing n free radicals (normalized so that No+ Nl = 1 and thus A = Nl), pA is entry rate coefficient in the absence of exit and a is the exited free-radical fate parameter.The fate parameter a is formally defined such that akfi is equal to the change in the entry rate coefficient arising from the Occurrence of exit. Thus its numerical value is restricted to the range - 1 < a 6 1. A value of a = + 1 would imply total re-entry of exited free radicals into the latex particles whereas a value of a = - 1 would imply that exited free radicals underwentWHANG, NAPPER, BALLARD, GILBERT A N D LICHTI 1119 complete cross-termination in the aqueous phase with entrant free radicals generated as a result of initiator decomposition. Values of a between these two extremes would imply that both re-entry and cross-termination in the aqueous phase could contribute to radical loss from the aqueous phase, as indeed could self-termination.More than one physical interpretation of such intermediate values for a is thus possible (see below). Eqn (1) and (2) may be derived by considering the kinetic equations governing the concentrations of aqueous-phase free radicals arising from (i) initiator decomposition and (ii) exit; these kinetic equations incorporate both self- and cross-termination of these species in the aqueous phase. Eqn (1) and (2) are then obtained by invoking the steady-state approximation for both types of aqueous-phase free radicals. Note that one physical interpretation of a = 0 is the absence of re-entry whence eqn (1) and (2) assume the form discussed previously. Eqn (2) can be re-written as: where f = -2ak, g = -2pA-(1 - a ) k , h = pA and we approximation in the aqueous phase.Integration of eqn (3) leads to the result where p - q6ert 1 -beyt Nl = (3) will adopt the steady-state (4) During interval 11, when the concentration of monomer is essentially constant, the fractional conversion x is given by t x = A J N,(t)dt where A = k, N, CM/nrNA, k, is the propagation rate constant, N, is the number of seed latex particles, CM is the monomer concentration in the particles, n f is the number of moles of monomer initially present and NA is Avogadro's constant. Eqn (5) shows clearly that at long times the time independent slope (a) of a plot of x against t is Aq so that q = B,,, where the subscript ss denotes the steady state. The intercept of this plot is given by b = lim [A(q--p)/y]ln {e;:-q - tdcn = [A(q -P)/Yl In (6/[6 - 11).(6) The expressions derived above permit a and b to be calculated given values for A , pA, k, a and Nl(t = 0). What is required, however, is to be able to calculate p A and1120 SEEDED EMULSION POLYMERIZATION k given values for a, b, A , a and Nl(t = 0). This inversion can be accomplished after some tedious algebra, the results being where k = A In F/2ab (7) P A = G k G = [2aa2+A(1 -a)a]/A(A--2a) (8) F = 0.5+{[2G+(l -a)+4aN1(t = 0)]/2[4G2+(a- 1)2+4G(a+ 1)]4). and Eqn (7) is clearly inapplicable if a = 0. The relationship k = [AN,(? = 0) - U] ( A - 2a)/Ab (9) must then be employed in conjunction with eqn (8) and the appropriate value of G.l RESULTS AND DISCUSSION R AD1 C AL RE-EN TRY The foregoing theoretical analysis enables kinetic data for the approach to the steady state to be analysed for different assumed values of a.Acceptable values for a can be ascertained by determining whether they predict realistic values for k and p A . Values of p A are found to be relatively insen'sitive to the assumed value for a and so provided the calculated value for p A does not exceed the experimentally determined rate of production of free radicals, p A is not a very discriminating parameter. Fortunately, at least at lower initiator concentrations, k is a discriminatory parameter, 150 50 100 150 so 100 time/min FIG. 1 .-Fractional conversion against time curves for seeded emulsion polymerizations: curves 1, 2 and 3 correspond to particle number concentrations of 3.0 x 10l6, 1.5 x 10'6 and 0.75 x 10l6 dm-3, respectively. Initiator concentrations: (a) 1.3 x 10-2, (b) 1.1 x 10-3, (c) 1.0 x 10-4 and (d) 1.2 x rnol13rn-~.WHANG, NAPPER, BALLARD, GILBERT A N D LICHTI 1121 being quite sensitive to the chosen value for a.Two criteria exist for the acceptability of k values: first, the values of k calculated at different particle number concentrations and different initiator concentrations must be identical, within experimental error; secondly, the absolute magnitude of this constant value for k measured for chemically initiated systems must agree with the value obtained by other independent methods (e.g. relaxation studies2). Fig. 1 displays the fractional conversion against time curves for seeded emulsion polymerizations of styrene.These were obtained at three particle number concentra- tions and at four different initiator concentrations lying in the range 10-5-10-2 mol dma3 potassium peroxydisulphate. The exit rate constant was measured independently by relaxation studies on a seeded emulsion polymerization of styrene initiated by prays2 The results were interpreted using the theory set out above. In relaxation studies, the entry rate is very small after removal of the initiating source. As a result, there are insufficient free radicals for cross-termination in the aqueous phase to occur (see below) and so the appropriate value of a was assumed to be + 1 (i.e. complete re-entry). The result so obtained, k = (2.6k0.4) x s-l, is in fair agreement with the value [( 1.4 f 0.4) x s-l] expected from previous measurements' on different latices. This latter value for k was computed using the size dependence TABLE CALCULATED VALUES OF pA AND k FOR DIFFERENT VALUES OF a pA/10-3 S-1 k/io-3 s-1 A,, a = + l a=O a = - 1 a = + l a=O a = - 1 111 N C /mol dm-3 / 1OI6 dm-3 1.2x lo+ 1.0 x 10-4 1.1 x 10-3 1.3 x 2.9 0.06 1.5 0.13 0.72 0.20 3.0 0.10 1.5 0.17 0.76 0.27 2.9 0.15 1.4 0.27 0.79 0.28 3.1 0.27 1.6 0.28 0.77 0.38 0.18 0.47 1.1 0.29 0.58 1 .o 0.9 1 1.2 2.4 2.1 3.1 4.0 0.25 0.26 0.63 0.67 1.4 1.5 0.39 0.41 0.76 0.81 1.3 1.4 1.2 1.3 1.5 1.7 2.9 3.2 2.6 2.8 3.7 4.0 3.5 4.9 mean standard deviation 20 10 12 8.1 6.6 3.3 3.9 6.9 6.7 8.2 3.6 8.6 4.9 14 3.5 3.6 4.1 3.1 2.9 2.2 5.6 2.6 4.7 4.4 5.6 3.0 3.8 1.1 1.9 2.2 2.8 1.8 1.9 1.6 3.5 1.9 3.6 3.3 4.3 2.6 2.5 1 .o k from relaxation studies = (2.6f0.4) x s-l.determined from previous studies. Some variation in the value of k would be expected for latices of the same size if the particles incorporate differing amounts of a soap that functions as a chain-transfer agent.' An analysis of the rate curves shown in fig. 1 for three different values of a, using the theory detailed above, is presented in table 1. The values of the average number of free radicals per particle in the steady state (ass) were all significantly less than 0.5. The values of k calculated for a = + 1 are clearly not constant. Moreover, the absolute values calculated for this value of a from these chemically initiated studies are significantly larger than the value of k measured for this latex at 50 *C using relaxation 37 FAR 11122 SEEDED EMULSION POLYMERIZATION kinetics. For both these reasons, the model with a = + 1, which corresponds to complete re-entry of exited free radicals, can be eliminated as a possible model describing these seeded styrene emulsion polymerizations in the presence of aqueous- phase initiator.Such systems are closely analogous to interval I1 of an ab initio polymerization. In contrast to the results obtained for a= + 1, the results for a = - 1 seem totally acceptable, both on the basis of the constancy of k and also in terms of its absolute magnitude. Indeed, the mean value obtained using a = - 1, k = (2.5+ 1.0) x s-l, is in remarkably good agreement with the relaxation value of k = (2.6 f 0.4) x s-l, although the large standard deviation associated with the former result suggests that such excellent agreement may well be fortuitous.The consistent results obtained with a = - 1 imply that for styrene, essentially complete cross-termination of the exited free radicals occurs in the aqueous phase by free radicals generated as a result of initiator decomposition. This would be in accord with the known8 rapid rate ofcross-termination of free radicals compared with the respective rates of self-termination. Typically, cross-termination rate constants are one or two orders of magnitude larger than those for self-termination. The enhanced rate of cross-termination may be a consequence of differences in the polarity of the two free-radical specie^,^ as well as the steric hindrance accompanying head-to-head linkage in self-termination.Note that the two free radical species in these studies are chemically quite different, the exited one being CH,=c-C6H, and the other SO,- or an oligomeric derivative thereof. Note that in the foregoing discussion we have assumed by analogy with copolymerization systems that low molecular weight free radicals cross-terminate more rapidly than they undergo self-termination. Unfortunately, although it is possible to exclude the model with a = + 1 (i.e. complete re-entry), it is more difficult to distinguish between values of cc = 0 and a = 1. The values of k calculated with a = 0 are reasonably constant and their mean value is k = (3.8 s-l; although large, this value is not totally incompatible with the relaxation result (2.6 x s-l).One simple interpretation of a = 0 would be that one half of the exited free radicals undergo cross-termination in the aqueous phase whilst the other half re-enter the latex particles. Other interpretations of a = 0 are, however, possible : e.g. the occurrence of complete self-termination in the aqueous phase of the exited free radicals, although this seems to be less likely on chemical An additional, and quite probable, interpretation of a = 0 is that the latex radical capture efficiency is low while cross-termination in the aqueous phase is very rapid. An exited free radical would then undergo essentially instantaneous cross- termination on entry into the aqueous phase which (because of the high rate of mutual-termination among aqueous-phase free radicals originating directly from the initiator) would have a negligible effect on p A .The foregoing results show that for the seeded emulsion polymerization of styrene in the presence of added initiator, complete re-entry of exited free radicals does not occur. The results suggest that significant cross-termination of the exited free radicals by free radicals generated by initiator decomposition takes place in the aqueous phase. 1.1) x DEPENDENCE OF THE RATE OF POLYMERIZATION O N THE NUMBER OF SEED LATEX PARTICLES The Smith-Ewart case 2 corresponds to Hss = 4. If case 2 kinetics are operative, the rate of polymerization would be directly proportional to the number concentration (E,,) of latex particles, as ass is independent of p, k and c.As van der Hoff has pointed out, the results of experiments are not always in accord with this simple model. Smith,loWHANG, NAPPER, BALLARD, GILBERT A N D LICHTI 1123 TABLE 2.-RATE OF POLYMERIZATION AS A FUNCTION OF THE NUMBER CONCENTRATION OF LATEX PARTICLES Nc tiss/ 1 0l6 dmd3 [I]/mol dm-3 NC/lOl6 dmW3 = 1 . 2 ~ 10-5 1 . 0 ~ 10-4 1.1 x 10-3 1.3 x 3.0k 0.1 0.18 0.30 0.44 0.84 1.5 f 0. I 0.20 0.25 0.39 0.44 0.75 & 0.05 0.14 0.21 0.22 0.29 C Q8 rate exponent vo. 2 j q . 2 q . 5 Bovey et aZ.ll and van der Hop2 have all reported experiments in which the rate was not directly proportional to the number of particles. Such results would be expected if A,, # f so that the value of A,, depends upon p, k and c ; in particular, p may then be a function of Nc so that the direct proportionality between rate and f l c would be lost.Table 2 presents values of Nc ass, which is proportional to the rate of polymerization, for different values of Nc over the range of initiator concentrations studied. Also presented in the table are values of the exponent y in the equation rate cc x[ obtained from the appropriate log-log plot. At low initiator concentrations when A,, is small, the exponent is only ca. 0.2. As A,, increases with the initiator concentration so does the exponent, reaching a value of ca. 0.8 for the highest initiator concentration studied. For the small particle size used in these experiments, it was not possible to achieve fi,, = 4; the initiator concentrations required to achieve this value of A ~ , induced particle coagulation.It should be stressed that the rate of polymerization of a system obeying Smith-Ewart kinetics is only directly proportional to the number of latex particles for case 2 kinetics (i.e. R,, = i), Deviations from an exponent of unity may well reflect deviations from the value of A,, = 4. At low initiator concentrations, the radical capture efficiency would be expected to approach 100% (see below). Under these conditions the total rate of entry of free radicals into latex particles Ncp would be constant, i.e. p a (RJ. Moreover, at low as,, fi,, = p/(2p+_k) - p / k , as k >> p.l Thus n,, would be inversely proportional to gc and the product NcnSs would be independent of Xc. Thus, for small nSs values and high radical capture efficiencies, the polymerization rate would be independent of the number of particles at constant initiator concentration.The exponent of Nc in the rate equation would then be zero. However, as R ~ , increases to 9, the exponent should increase towards unity. This is qualitatively what was observed in these studies. We conclude that deviations from unity in the exponent of xc do not necessarily reflect deviations from the general Smith-Ewart-type kinetics, merely deviations from Smith-Ewart case 2 kinetics. These deviations arise in part from variations in the radical capture efficiency and thus in p. Smithlo originally proposed that these deviations were a consequence of the inability of monomer to diffuse sufficiently rapidly from the droplets to the polymerizing particles to sustain propagation.This explanation seems unlikely in light of Flory's calc~lation~~ that monomer diffusion is comparatively rapid in these systems. 37-21124 SEEDED EMULSION POLYMERIZATION TABLE 3 .--FREE-RADICAL ENTRY RATES AT DIFFERENT PARTICLE NUMBER CONCENTRATIONS (a= -1) [I]/mol dm-3 = Nc/ 1 0l6 dmP3 1 . 2 x 10-5 1.0 x 10-4 1.1 x 10-3 1.3 x 3.0 f 0.1 0 . 1 4 0.23 0.78 1 . 9 1.5 f 0.1 0.38 0 . 4 9 1 . 2 2.8 0.75 k 0.05 0 . 9 4 0.97 2.2 3.9 TABLE 4.-RADICAL CAPTURE EFFICIENCIES AS A FUNCTION OF PARTICLE NUMBER CONCENTRATION (a= - 1) NebA - knss)/ l OI3 dm-3 s-l capture efficiency (%) IIl/mol RC/1Ol6 dm-3 = dm-3 1.2 x 10-3 1.0 x 10-4 1 . 1 x 10-3 1.3 x 10-2 1.2 x 10-5 1.0 x 10-4 1 . 1 x 10-3 1.3 x 10-2 3.0k0.1 0.41 0.69 2.3 5.9 21 3 1.2 0.3 1.5 +_ 0.1 0.58 0.73 1.7 4.5 29 3 0.9 0.2 0.75 +_ 0.05 0.68 0.74 1.7 3.0 34 3 0.9 0.2 RADICAL CAPTURE EFFICIENCIES In what follows, we will adopt the value a = - 1.As noted above, however, the calculated values of p A are relatively insensitive to the assumed value of a (see table 1). Table 3 displays the values of the average rate of free radical entry per particle (PA - kfi,,) at different particle number concentrations. Lower particle number concentrations could not be investigated because of the occurrence of nucleation in the seeded systems. It is clear that the larger the particle number concentration, the smaller is the rate of free radical entry per particle. Table 4 shows that the overall rate of entry of free radicals into the latex particles, Nc(PA - kfiss), is relatively insensitive to the particle number concentrations at a given initiator concentration.The overall radical capture efficiency, calculated as described previously,' is a sensitive function of the initiator concentration; efficiencies varied in these experiments from 0.3 to 21% as the initiator concentration decreased from 1.3 x mol dm-3. The efficiency was relatively insensitive to the particle number concentration over the four-fold range of concentration studied. This observation is in accord with the predictions of the theory of Hawkett et aZ.14 Briefly, since most of the bimolecular termination of free radicals generated by initiator decomposition occurs in the aqueous phase, the presence of the latex particles has virtually no influence on the radical capture efficiency, at least over the range of particle concentrations accessible in these studies.to 1.2 xWHANG, NAPPER, BALLARD, GILBERT A N D LICHTI 1125 CONCLUSIONS The results of these studies on the seeded emulsion polymerization of styrene rule out the occurrence of complete re-entry of exited free radicals. The data are best fitted by a model that assumes almost complete cross-termination of exited free radicals in the aqueous phase by the free radicals generated as a result of decomposition of the initiator. This is in accord with the known rapidity of cross-termination reactions compared with the corresponding self-termination reactions. It is also shown that the average rate of entry of free radicals into individual latex particles increases with decreasing particle number concentration at constant initiator concentration.The overall radical capture efficiency was, however, relatively insensitive to the particle number concentration over the range studied. These results may be quite simply interpreted as arising from moderately extensive termination taking place in the aqueous phase. Finally, it was found that the rate of polymerization was a complex function of the number concentration of latex particles. This is because this rate is, from the general Smith-Ewart kinetic scheme presented in this paper, seen to be a known but fairly complex function of quantities such as p A , a, etc. Each of these latter are simple functions of Nc, or at least can be represented in terms of straightforward kinetic equations involving various types of aqueous-phase free radicals.Thus the behaviour of the rate of polymerization as Nc is varied reflects deviations from Smith-Ewart case 2 kinetics (nss = i), rather than deviations from the general Smith-Ewart kinetic scheme or the onset of diffusion control of the propagation step. LIST O F SYMBOLS a b C 2 k R k n, = slope of steady-state rate plot (s-l) = intercept of steady-state rate plot = first-order (with respect to particles) bimolecular termination rate constant = concentration of monomer in the latex particles (mol dm-3) = initiator concentration (mol dm-3) = exit rate coefficient (s-l) = propagation rate constant (dm3 mol-1 s-l) = number of moles of monomer initially present = average number of free radicals per particle = steady-state value of A = Avogadro’s constant (mol-l) = relative population of particles containing n free radicals = total number of seed latex particles = number concentration of seed latex particles (dm-3) = time (s) = exited free-radical fate parameter (- 1 < a < 1) = free-radical entry rate coefficient (s-l) = free-radical entry rate coefficient in the absence of exit (s-l).(s-Y We thank the ARGC for financial support of these studies and AINSE for a post-doctoral fellowship for G. L. ; M. J. B. acknowledges the award of a Sydney University post-graduate scholarship. We thank the Electron Microscope Unit of the University of Sydney for their generous provision of facilities.1126 SEEDED EMULSION POLYMERIZATION B. S. Hawkett, D. H. Napper and R. G. Gilbert, J. Chem. Soc., Faraduy Trans. I , 1980, 76, 1323. S. W. Lansdowne, R. G. Gilbert, D. H. Napper and D. F. Sangster, J. Chem. SOC., Faraduy Trans. I , 1980,76, 1344. J. Ugelstad and J. K. Hansen, Rubber Chem. Technol., 1976, 49, 536. J. W. Vanderhoff, in Vinyl Polymerization, ed. G. E. Hamm (Marcel Dekker, New York, 1969), vol, 1, part 2, chap. 1. D. C. Blackley, Emulsion Polymerization (Applied Science, London, 1975). W. V. Smith and R. H. Ewart, J. Chem. Phys., 1948, 16, 592. D. T. Birtwistle and D. C. Blackley, J. Chem. SOC., Faraday Trans. 1, 1981,77, 413. H-G. Elias, Macromolecules (Plenum, New York, 1977), vol. 2, p. 7. 1979). * M. J. Bowden, in Macromolecules, ed. F. A. Bovey and F. H. Winslow (Academic Press, New York, lo W. V. Smith, J. Am. Chem. Soc., 1948, 70, 3695. I1 F. A. Bovey, I. M. Kolthoff, A. J. Medalia and E. J. Meehan, Emulsion Polymerization (Wiley, New l2 B. M. E. van der Hoff, J. Phys. Chem., 1956, 60, 1250. l3 P. J. Flory, Principles of Polymer Chemistry (Cornell University Press, Ithaca, 1953), p. 210. l4 B. S. Hawkett, D. H. Napper and R. G. Gilbert, J. Polym. Sci., Polym. Lett. Ed., submitted for York, 1955). publication. (PAPER 1/727)

ISSN:0300-9599    

DOI:10.1039/F19827801117

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1982, 78, 1127-1 130 Magneto-optical Rotation Studies of Liquid Mixtures Part 5.-Binary Mixtures of Water with Dimethylsulphoxide, Formamide, Methyl Formamide, Dimethyl Formamide and Acetonitrile BY J. GRAHAM DAWBER Department of Chemistry and Biology, North Staffordshire Polytechnic, College Road, Stoke-on-Trent ST4 2DE Received 7th May, 1981 Measurements have been made of the magneto-optical rotations at 435.8 nm of binary liquid mixtures of water with dimethylsulphoxide, formamide, methyl formamide, dimethyl formamide and acetonitrile as part of a study to investigate the applicability of the magneto-optical method in detecting weak molecular interactions in liquid mixtures. The departure from ideality of the mixtures was treated in terms of an excess magnetic rotation, aE.For mixtures of water with dimethylsulphoxide, formamide, methyl formamide and dimethyl formamide the values of aE indicated interaction between the components, whereas with acetonitrile the aE values were compatible with a breakdown of the water structure in the mixtures. In previous s t ~ d i e s l - ~ of the magneto-optical rotation of binary mixtures of miscible liquids it was shown that the departures from ideality of the systems could be expressed in terms of an excess magnetic rotation, aE. It was found that the variation of aE with liquid composition could be correlated with changes in the structural properties of the liquid mixtures as observed by other workers using different techniques. In the work presented here the previous have been extended to binary systems of water with dimethylsulphoxide (DMSO), formamide, methyl formamide (MF), NN-dimethyl formamide (DMF) and acetonitrile in order to test further the applic- ability of the magneto-optical method in detecting changes in constitution in the liquid state.EXPERIMENTAL The magnetic rotation apparatus has been described previously.'* 6 v The organic liquids were dried with anhydrous MgSO, and distilled prior to use. Binary mixtures of water plus the second component were made up to cover the whole composition range. Measurements of magnetic rotation were made on each solution, nominally at 20 O C , using a 1 cm path-length cell placed in a magnetic field of flux density of 0.54 T, and using the 435.8 nm line of the Hg emission spectrum. Readings of angular rotation could be estimated to O.O0lo.RESULTS The results were quoted as magnetic rotation, a,, defined as a, = O,/l (1) where Os is the magnetic rotation of the solution in a magnetic flux density of 1 T and 1 is the path-length of the solution in metres. The units of a, are thus O T-l m-l. The values of magnetic rotation for the single liquids, a:, are given in table 1. 11271128 MAGNETO-OPTICAL ROTATION STUDIES OF LIQUID MIXTURES TABLE 1 .-MAGNETO-OPTICAL ROTATIONS OF PURE LIQUIDS (a:) AT 436 nm AND 20 O C liquid a",/" T-l m-l water DMSO acetonitrile formamide N-methyl formamide NN-dimethyl formamide 392.8 639.3 308.9 468.9 435.0 402.1 FIG. 1.-Excess magnetic rotations (aE) ,of binary mixtures with water: (a) DMSO; (b) acetonitrile; (c) formamide; (d) methyl formamide; (e) dimethyl formamide.J.G. DAWBER 1129 The non-ideality of the liquid mixtures was expressed in terms of the excess magnetic rotation, aE, which is defined as a" = a,-(X,aO,,+X,aO,,) where X, and X z are the mole fractions of components 1 and 2 in the mixture and agl and a;, are the magnetic rotations of the pure components. The values of aE for the various binary liquid systems are plotted as a function of XHzO in fig. 1. There are several possible reasons for the changes in aE as the liquid composition is varied. These include association of the components or breakdown of the liquid structure of one or both of the components, or, alternatively, specific molecular interaction between the components.From previous it was found that negative values of aE could be correlated with breakdown of the liquid structure of one component by the other and that positive values of aE could be correlated with interaction between the components. DISCUSSION Many different studies have been made of the physical properties of the DMSO + water system in order to investigate the molecular structure and motion of the liquid as a function of compo~ition.~-~~ All of these studies indicate a strong interaction between DMSO and water with a maximum in the interaction at a mole fraction of water, XHzo, between 0.6 to 0.7. Inspection of the aE curve for the DMSO + water system (fig. 1) indicates a large deviation from ideality. From previous this large positive aE effect indicates a strong specific interaction between the components, likely to be hydrogen bonding, which is at a maximum at XHz0 = 0.7; this is in complete agreement with the studies of other workers using different techniques.The acetonitrile+water system has also been studied by many different tech- n i q u e ~ . ~ ~ - ~ ~ These studies in general show that the components of the system do not participate in specific interactions, but rather that the acetonitrile causes a break- down of the water-water interactions. The aE values from this work for the acetonitrile + water system (fig. 1) are negative over the whole range of concentration and, by analogy with previous work,' this would indicate structure-breaking of the water by the acetonitrile with no specific interaction between the two components; this is supported by most studies involving other technique^.^^-^^ The individual components in the formamide + water system are themselves capable of forming three-dimensional hydrogen-bonded associated aggregate^.,^' 26 While mixtures of formamide and water cannot be regarded as ideal, several s t ~ d i e s ~ ~ - ~ ~ suggest that the deviations from ideality are small but that, nevertheless, a formamide- water complex is formed in the mixtures.The high compatibility of the two solvents could be due to the intermolecular hydrogen-bond energy of the formamide-water complex being similar to the values for the pure solvents themsel~es.~~ Despite the apparent ideality of formamide+water mixtures, the mixed system is a complex medium in which interaction between the two species does occur. This is apparent from the aE curve (fig.1) for the system, which exhibits positive values over the whole range of composition with a maximum at XHzo = 0.5. The symmetric nature of the aE curve is compatible with the formation of a 1 : 1 formamide-water complex. The effect for formamide, however, is much smaller than for DMSO. The aE curves for the MF and DMF systems with water (fig. 1) show similar deviations from ideality indicating a maximum in the curves at XHzo x 0.7. Measure- ments of viscosity in the DMF+water system exhibit a maximum at XH,o x 0.751130 and the complex DMF-3H20 has been ~ u g g e s t e d . ~ ~ ? ~ ~ The aE results of this work support this proposal, with a similar, but marginally smaller, interaction for MF as well as for DMF.Thus, although the magneto-optical technique is not one of direct observation of complex formation, in that it assumes that departures from an ideal mixture law can be interpreted in terms of molecular interactions, it appears that the a values correlate well with other observations on the behaviour of binary mixtures aimed at detecting weak molecular interactions between the components. MAGNETO-OPTICAL ROTATION STUDIES OF LIQUID MIXTURES The author gratefully acknowledges the financial support of the S.R.C. towards the cost of the polarimeter. J. G. Dawber, J. Chem. SOC., Faraday Trans. I , 1978, 74, 1702. J. G. Dawber, J. Chem. SOC., Faraday Trans. I , 1978, 74, 1709. J. G. Dawber, J. Chem. SOC., Faraday Trans. 1, 1979, 75, 370.J. G. Dawber, J. Chem. SOC., Faraday Trans. 2, 1980, 76, 1324. J. G. Dawber, J. Chem. SOC., Faraday Trans. I , 1978, 74, 960. J. G. Dawber, J. Chem. SOC., Faraday Trans. 2, 1974, 70, 597. 'I J. Kenttamaa and J. J. Lindberg, Suom. Kemistil., 1960, 33, 32 and 98. a M. E. Fox and K. P. Whittingham, J. Chem. SOC., Faraday Trans. I , 1975,75, 1407. lo J. M. G. Cowie and P. M. Toporowski, Can. J. Chem., 1961, 39, 2240. l1 G. J. Safford, P. C. Schaffer, P. S. hung, G. F. Doebbler, G. W. Bradyand E. F. X. Lyden, J. Chem. l2 G. Brink and M. Falk, J. Mol. Struct., 1970, 5, 27. l3 K. J. Packer and D. J. Tomlinson, Trans. Faraday SOC., 1971,67, 1302. l4 S. Y. Lam and R. L. Benoit, Can. J. Chem., 1974,52, 718. l5 J. C. Chan and A. W. van Hook, J. Solution Chem., 1976, 5, 107.l6 0. Kiyohara, G. Perron and J. E. Desnoyers, Can. J. Chem., 1975, 53, 3263. l7 G. E. Walrafen, J. Chem. Phys., 1970, 52,4176. D. H. Rasmussen and A. P. Mackenzie, Nature (London), 1968, 220, 1315. l9 B. G. Cox, R. Natarajan and W. E. Waghorne, J. Chem. SOC., Faraday Trans. I , 1979,75, 86. 2o C. de Visser, W. J. M. Heuvesland, L. A. Dunn and G. Somsen, J. Chem. SOC., Faraday Trans. I , 21 K. W. Morcom and R. W. Smith, J. Chem. Thermodyn., 1969, 1, 503. 22 D. A. Armitage, M. J. Blandamer, M. J. Foster, N. J. Hidden, K. W. Morcom, M. C. R. Symons 23 G. Wada and S. Umeda, Bull. Chem. SOC. Jpn, 1962, 35, 1797. 24 E. v. Goldammer and H. G. Hertz, J. Phys. Chem., 1970, 74, 3734. 25 J. M. McDowall and C. A. Vincent, J. Chem. SOC., Faraday Trans. I , 1974, 70, 1862. 26 J. M. McDowall, N. Martinus and C. A. Vincent, J. Chem. Soc., Faraday Trans. 2, 1976, 72, 654. 27 C. D. Sinclair and C. A. Vincent, J. Chem. SOC., Faraday Trans. 1, 1974,70, 1926. 28 P. Rhodewald and M. Moldner, J. Phys. Chem., 1973, 77, 373. 29 A. Fratiello, Mol. Phys., 1963, 7, 565. 30 A. Johannson and P. A. Kollman, J. Am. Chem. SOC., 1972,94,6196. 31 D. Singh, L. Bahadur and M. V. Ramanamurti, J. Solution Chem., 1977,6, 703. 32 L. Bahadur and M. V. Ramanamurti, J. Chem. SOC., Faraday Trans. I , 1980, 76, 1409. H. L. Clever and S. P. Pigott, J. Chem. Thermodyn., 1971, 3, 221. Phys., 1969,50, 2140. 1978, 74, 11 59. and M. J. Wooten, Trans. Faraday SOC., 1968, 64, 1193. (PAPER 1 /728)

ISSN:0300-9599    

DOI:10.1039/F19827801127

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1982, 78, 1131-1 140 Change in the Phase Stability of Zinc Blende and Wurtzite on Grinding BY KENICHIMAMURA AND MAMORU SENNA* Department of Applied Chemistry, Faculty of Engineering, Keio University, 3-14-1 Hiyoshi, Yokohama 223, Japan Received 1 lth May, 1981 Two different zinc blende (ZB) powders were vibro-milled and the transformation process into the wurtzite (WZ) phase on subsequent heating in N, was observed. Vibro-milled ZB transformed rapidly into WZ at the beginning of the heating. Part of the WZ was then retransformed into ZB leaving a large amount of WZ, the reaction attaining an apparent equilibrium state. The transition took place even at 1023 K, which is more than 200 K lower than the lowest Z B e WZ transition temperature ever reported.The results were interpreted on the assumption of an increase in the chemical potential and a broadening of the chemical potential distribution within the powder mass caused by the mechanical treatment. Unlike liquids and gaseous substances, the chemical potential of a crystalline solid is generally structure sensitive. This often leads to uncertainty about the conditions necessary for phase stability. Reported values of the transition temperature between zinc blende and wurtzite, for example, are widely scattered, i.e. 1230,' 1243,2 1248,3 1293* or even 1423 K5 The variation in stability region caused by different degrees of activity will become more marked when a solid is mechanically treated. This is one of the fundamental principles of mechanochemical activation.6 There are very few systematic studies, however, on the mechanochemical effect on the polymorphic equilibrium.Zinc sulphide is a suitable material for a systematic study because of its high structural sensitivity. Gmk7 studied the mechanochemical activation of ZnS with regard to ore dressing. In his work, however, phase transformation played a subsidiary role. The main purpose of the present study is to elucidate the change in phase relation between zinc blende and wurtzite caused by preliminary grinding. EXPERIMENTAL MATERIALS Two different sorts of zinc blende were used as starting materials. (a) Synthetic ZnS8 (ZP): gaseous H,S was blown slowly into 0.1 mol dm-3 ZnSO, aqueous solution for 10 h. Precipitated ZnS was washed with water and acetone and was dried at room temperature under reduced pressure for 2 days.After deagglomeration in a mortar, the sample was tempered in N, at 673 K for 2 h and then at 753 K for 3 h. The temper conditions were chosen so as to obtain as good crystallinity as possible with no transformation into wurtzite. (b) Natural ZnS (ZN): A natural zinc blende ore from Aniya, Japan, was crushed under 100 mesh and tempered in N, at 473 K for 5 h. 11311132 PHASE STABILITY OF ZnS GRINDING The starting material was vibro-milled in an inert liquid, cyclohexane, in order to reduce the possible complicating effects of the moisture, oxygen and significant local heating. A 1 g sample, 100 steel balls of 6 mm diameter and 30 cm3 cyclohexane were put into a 50 cm3 cylindrical steel vessel of 35 mm diameter. Vibro-milling was carried out isothermally at 303 K using a small laboratory mill (Glen Creston) operated at 12 Hz.After grinding, the supernatant was decanted and the rest was dried for 15 h at room temperature under reduced pressure. The specimens are labelled ZP-n or ZN-n according to the starting material, with n being the grinding time in hours. HEATING OF THE SPECIMEN The samples were heated at predetermined temperatures iegulated within 1 K in a half-opened tube either of glass (< 873 K) or of quartz (2 873K) under a N, flow of 5 dm3 min-l. X-RAY DIFFRACTOMETRY Because of the superimposition of many diffraction peaks, it is not easy to determine the amount of each phase precisely for the 11-VI compounds which have zinc blende and wurtzite isomorphic ~tructures.~ The difference in the interplannar distance between zinc blende (1 1 1) and wurtzite (002) is only 5 x nm.Hence the diffraction peaks from these lattice planes could not be resolved. Instead, the total intensity of these two peaks, I,, was determined. On the other hand, the ratio of the intensity of wurtzite (loo), I , (loo), to I, was constant at 1.3 for all the samples obtained by heating at > 1273 K for a sufficiently long time. By assuming that the ratio I,(lOO)/~, = 1.3 corresponds to the pure wurtzite phase, where 1,(111) = 0 [cJ the ratio I,( 100)/1,,,(002) = 1.16, according to the ASTM card], the fraction of wurtzite phase, (1) x,, was determined by X, = 1,(100)/1.31I, where I, = Z,(002) + I,( 1 1 1). The lattice distortion and the crystalline size were determined by conventional methods according to Hall.lo RESULTS PHASE TRANSFORMATION OF GROUND ZnS A remarkable phase change process was observed on heating the ground ZnS isothermally in N, flow. As shown in fig. 1, there was a rapid phase transformation into wurtzite at an early stage of heating, showing a maximum wurtzite fraction, xwm. A sluggish retransformation then occurred, resulting finally in an apparent equilibrium state at the fraction of wurtzite, x,,. Note that the transformation was scarcely observed on very well crystallized zinc blende, ZN-0, even after heating at 1273 K for 3 h. This indicates that the initial rapid transformation into wurtzite is characteristic of ground zinc blende. With increasing heating temperature, both x, and x,, increased as shown in fig.2 and 3. However, the rate of increase with temperature was much larger for the ZP series than for the ZN series. The effect of grinding, particularly on x, was also much more pronounced for ZP than ZN. RELATION BETWEEN X, X, A N D THE LATTICE DEFORMATION The distortion of the zinc blende lattice, q-, increased, whereas the crystallite size, D, decreased with grinding time in the manner shown in fig. 4. Before heating, no mechanochemical phase transformation into wurtzite was observed. Within each series (ZP or ZN), x, increased monotonically with q, as shown in fig. 5 . The value of x, at the same q, however, was substantially higher for ZP than ZN. The relationK. IMAMURA AND M. SENNA 1133 0.71 0 10 20 30 heating time/h FIG.1.-Change in the fraction of wurtzite, x,, with heating time for ZN-8 (squares) and ZP-8 (circles). Curves 1 and 4, heated at 1023 K; curves 2 and 5, heated at 1123 K; curves 3 and 6, heated at 1173 K. 0.7 0.6 0.5 0 . 4 E Y 0.3 0.2 0 . 1 f' 1 0 1 " ' I " " I " " 1 ~ ' 1050 1100 11 50 temperature/K FIG. 2.-Relation between x, and temperature. Curve 1, ZN-4; curve 2, ZN-8; curve 3, ZP-0; curve 4, ZP-4; curve 5, ZP-8.1134 PHASE STABILITY OF ZnS 0.4 - 083 - -.---------:: 1 0.1 - -0 I , , I , , , I , , 0 1050 1100 1150 temperature/K FIG. 3.-Relation between x, and temperature. Curve 1, ZN-4; curve 2, ZN-8; curve 3, Zp-0; curve 4, ZP-4; curve 5, ZP-8. 3 0 1 2 3 4 5 6 7 8 melting time/h FIG. 4.-Variation of lattice distortion (filled symbols) and crystallite size (open symbols) with milling time, for ZN (0, m) and ZP (0, a).K.IMAMURA A N D M. SENNA 1135 0.7 0 . 6 0.5 0 . L E 2 0.3 0.2 0 . 1 lattice distortion.,'q (arb. units) FIG. 5.-Relation between x, and lattice distortion for ZN (squares) and ZP (circles). Curves 1 and 4, heated at 1023 K; curves 2 and 5, heated at 1123 K; curves 3 and 6, heated at 1173 K. 1 .o 0.5 5 0.2 0 . 1 8 10 20 50 80 100 crystallite size, D/nm FIG. 6.-Relation between x, and crystallite size. Curve 1, heated at 1023 K; curve 2, heated at 1123 K; curve 3, heated at 1173 K.1136 PHASE STABILITY OF ZnS FIG. 7.-Relation between x,, and lattice distortion for ZN (squares) and ZP (circles). Curves 1 and 4, heated at 1023 K; curves 2 and 5 , heated at 1123 K; curves 3 and 6, heated at 1173 K.8 10 20 50 80 100 crystallite size, D/nm FIG. 8.-Relation between x,, and crystallite size. Curve 1, heated at 1023 K ; curve 2, heated at 1123 K; curve 3, heated at 1173 K. between xwm and D, on the other hand, was fairly simple, irrespective of the starting material, as shown in fig. 6. The relations between x,, and v (fig. 7) and x,, and D (fig. 8) were similar to those for x, with the exception that x,, was practically constant when q was large. DISCUSSION EFFECTS OF A N D D O N THE RATE OF TRANSFORMATION As seen from fig. 5-8, the effects of and D are more pronounced on x, than on x,,, because the imperfection of the matrix phase could serve as a driving forceK. IMAMURA A N D M. SENNA 1137 for the zinc blende + wurtzite transformation, by enhancing the nucleation of the new phase, but not for the reverse process.The distorted zinc blende will recover to a less distorted state, which is recognized from the sharpening of the X-ray diffraction peak of the starting phase, as shown in fig. 9.* The major portion of the decrease in the line breadth occurs at the initial stage of the heating. The behaviour is compatible - 0 10 20 30 40 heating time/h FIG. 9.-Variation of X-ray diffraction peak breadth of zinc blende ( 1 1 1 ) with heating time, for ZP-8. Curve 1, heated at 1023 K ; curve 2, heated at 1123 K; curve 3, heated at 1173 K. with the observation represented in fig 1, in the sense that the transformation from zinc blende to wurtzite occurs mainly at the early stage of the heating.It indicates that the transformation from zinc blende to wurtzite characteristic of the ground zinc blende is achieved at the cost of the lattice imperfection of the matrix zinc blende phase. On the other hand, x, and x,, were larger for smaller values of D at the same values of q. This would mainly be owing to the larger sites for the formation of the new stacking, i.e. the nucleation of the wurtzite phase. This was similar for the aragonite -+ calcite transformation,ll in spite of the rather different transformation mechanism. CHANGE I N THE TRANSFORMATION TEMPERATURE OWING TO G R I N D I N G The extent of the maximum transformation, x, is determined not only by the balance of the rates of the forward and reverse transformations, but also by the stability of the phases.As already mentioned, the reported temperatures of zinc blende G= wurtzite equilibrium range from 1230 to 1243 K. In the present experiment, the transformation of zinc blende into wurtzite occurred ,at 1023 K, which is more than 200 K lower than the lowest equilibrium temperature reported previously. The formation of intermediate phases possibly assigned as y-ZnS l2 or polytypes such as 6H, 18H or 24H l3 was not observed in the present experiment. Even if the transformation takes place through such intermediate structures, it still seems necessary to postulate that the chemical potential of the zinc blende phase has been changed significantly because of the mechanical treatment, as will be discussed later. * Because of the peak superposition, a precise analysis to obtain q and D separately for the partially transformed samples was not possible.1138 PHASE STABILITY OF ZnS POSSIBILITY OF THE GRADUAL TRANSITION MECHANISM Besides the maximum degree of transformation, one of the most characteristic observations was the existence of the apparent equilibrium state of the mixed phases, even after prolonged heating.According to classical thermodynamics, coexistence of two immiscible equicomponent phases at a constant pressure and an arbitrary temperature is not allowed. If two phases were miscible, forming a mixed stacking due to the layer structure of ZnS, then a temperature region for the stable coexistence of both phases could still be expected, by assuming the gradual transition mechanism proposed by Allen and Eagles.14 Even if the gradual transition mechanism were working and the different value of x,, at different temperature were explained, it still leaves open the question of why the retransformation from wurtzite to zinc blende takes place.NECESSARY ASSUMPTIONS In order to explain the above behaviour with regard to the phase transformation, the following two assumptions are necessary: (i) the chemical potential of the zinc blende phase, pZ, is increased on grinding and (ii) the value of pz is not identical throughout the sample but has a certain distribution within a powder mass, the distribution being broadened when the sample is ground. The major part of the excess free energy in active solids is considered to be the enthalpy contribution.15 Since a number of experimental r e s u l t ~ l ~ - ~ ~ verify the increase in enthalpy through mechanical stressing, assumption (i) is generally acceptable, A direct measurement of the enthalpy increase of the present specimen using a differential scanning calorimeter was not successful, presumably because of the insufficient sensitivity of the instrument and the sluggish energy release.Assumption (ii) can also be accepted, since each particle has a different chance of being hit by the milling balls. Even if the chance could be averaged after prolonged grinding, local differences in the microdeformation within particles cannot be avoided. Moreover, even the particle size distribution alone could result in the free energy distribution since, according to Allen and Eagles,l* the chemical potential of the solid material is a function not only of the temperature, pressure and composition, but also of the lattice distortion and the particle size.If the liberation of excess energy were easily observable, e.g. by means of d.t.a. or d.s.c. thermograms, as with A1,20 CaFZ2l or y-Fe20,,22 it would be possible to estimate the broadness of the distribution. With these assumptions, the behaviour of zinc blende during vibro-milling and subsequent heating can be explained. VARIATION OF THE CHEMICAL POTENTIAL DISTRIBUTION The chemical potential distribution, m), of well crystallized zinc blende, mZ)O, is very sharp, as shown schematically in fig. lO(a). No phase transformation into wurtzite can take place at a temperature, qx, below the ‘true’ transformation point, cr, where the two chemical potential curves, & and ,u:, of zinc blende and wurtzite in well crystallized states, respectively, cross each other.After grinding, the chemical potential distribution is broadened between pz and pz,max, as shown in fig. lO(b). Since a part of the ground sample has a chemical potential higher than &, it is possible for the transformation to take place at temperatures between ItJr,rnin and cr, as shown in fig. lO(c). The resulting wurtzite phase could also possess a certain distribution of chemical potential owing to the insufficient growth of the new nuclei. After prolonged heating, the rest of the zinc blende, which has not transformed into wurtzite, recovers to make the distribution narrower again, so that no further transformation into wurtzite is possible, as shown in fig.lO(d). On the other hand,T P ( d ) T T FIG. 10.-Schematic representation of changes in chemical potentials and their distribution caused by grinding and subsequent heating. (a) Starting material (well crystallized zinc blende); (b) ground zinc blende; (c) after heating for a short time; ( d ) after prolonged heating.1140 PHASE STABILITY OF ZnS some insufficiently grown wurtzite has a chance to retransform into zinc blende. The rest of the wurtzite grows further resulting in well crystallized wurtzite, as is also seen from the sharpening of the X-ray diffraction peaks (not shown). Thermodynamically, it is still possible for wurtzite to transform into zinc blende, even after the phase has grown, when heated at temperatures lower than TFr.In the actual experiments this did not occur. Instead, an apparent equilibrium state was observed. This could be due to kinetic factors, and in particular the difficulty of nucleation, as in the case of the massicot to litharge transformation without mechanical aids.23 COMPETITIVE PROCESSES The excess free energy in the activated zinc blende could be used up as a driving force either in the transformation to wurtzite, or in the return to the more stable zinc blende. through recovery and recrystallization. A similar kind of competition should also take place in the latter part of the transformation process, where the insufficiently crystallized wurtzite phase could either retransform to zinc blende or remain as wurtzite and continue to grow.Which of the two processes is dominant is an important question when mechanochemical activation is applied to practical solid-state reactions. It depends on many complicated factors, one of the most important being the heating temperature, as partly elucidated in the case of mechanically treated y-Fe,0,.22 We thank Prof. H. Kuno for valuable discussions. A. Kremheller and A. Lavine, Sylvania Technol., 1957, 10, 67. H. Samelson, J. Appl. Phys., 1962, 33, 1779. N. N. Sirota and V. P. Sapelkina, Krist. Tech., 1971, 6, 381. E. T. Allen and J. L. Grenshaw, 2. Anorg. A&. Chem., 1913, 79, 2. A. Addamino and M. Aven, J. Appl. Phys., 1960, 31, 1. * R. Schrader and B. Hoffman, Festkorperchemie (VEB Deutscher Verlag, Leipzig, 1973), p. 522. ' E. Gock, Habilitationsschrijit (T. U. Berlin, 1977). B. Basak, D. R. Glasson and S. A. A. Jayaweera, Particle Growth in Suspensions, SOC. Chem. Ind. Mongr., 1974, 38, 143. !a G. Ohtani and M. Senna, Phys. Status Solidi A, 1980, 60, K35. lo W. H. Hall, Proc. Phys. Soc. London, Sect. A , 1949, 62, 741. l1 H. Momota, M. Senna and M. Takagi, J . Chem. Soc., Faraday Trans. I , 1980, 76, 790. l2 D. C. Buck and L. W. Strock, Am. Mineral., 1955, 40, 192. l3 C. J. Schneer, Bull. Geol. SOC. Am., 1958, 69, 1640. l4 J. W. Allen and D. M. Eagles, Physica, 1960, 26, 492. l5 K. Torkar, Proc. 4th Int. Symp. Reactivity of Solids, Amsterdam, 1960, ed. J. H. deBoer (Elsevier, lo G. S. Chodakov and L. I. Edelmann, Kolloidn. Zh., 1967, 29, 728. Amsterdam, 1961), p. 400. R. Schrader and B. Hoffmann, 2. Anorg. Allg. Chem., 1969,369,41. K. TkaEova, Silikaty, 1976, 4, 321. M. Senna and K. Schonert, Powder Technol., to be published. 2o L. M. Clarebrough, M. E. Hargreaves, M. H. Loretto and G. W. West, Acta Metall., 1960, 8. 797. 21 R. Schrader and W. Oese, J. Therm. Anal., 1970, 2, 349. 22 H. Imai and M. Senna, J. Appl. Phys., 1978, 49, 4433. 23 M. Senna and H. Kuno, J. Am. Ceram. Soc., 1971,54, 259. (PAPER 1/752)

ISSN:0300-9599    

DOI:10.1039/F19827801131

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC. Faraday Trans. I, 1982, 78, 1141-1 148 Kinetics of the Gas-phase Addition Reactions of Trichlorosilyl Radicals Part 3.-Additions to 2-Olefins BY TAKAAKI DOHMARU* AND YOSHIO NAGATA Radiation Centre of Osaka Prefecture, Shinke-cho, Sakai, Osaka, Japan Received 1 1 th May, 198 1 The following Arrhenius parameters for the forward and reverse steps of trichlorosilyl radical additions to trans-but-2-ene, cis-but-2-ene, cis-pent-2-ene, 2-methyl-but-2-ene and cyclopentene have been obtained by a competitive method. The relevant elementary reactions are (3) * SiC1, + CH,COCH, + (CH,),cOSiCl, SIC^,+ +< >-$--s~cI, (5, - 5 ) \ I I I and c-y -SiCl, + HSiC1, + H~-~-SiCl, + SiC1, (6) / E5-E3 log(A-,/A, E+-EB olefin log (A,/A,) /kJ mol-l /mol ern-,) /kJ rno1-I temp./K cis-CH,CH=CHCH, 0.91+_0.02 0 4.09 k 0.10 77.0 f 0.8 397-462 trans-CH,CH=CHCH, 0.81+0.01 0 3.99 & 0.17 76.1 _+ 1.3 378-501 cis-CH,CH=CHCH,CH, 0.57+0.03 0 5.39 0.10 85.8 _+ 0.8 406-455 * * cis-CH,CH=CHCH,CH, 0.89+0.01 0 5.39k0.10 85.8k0.8 406-455 * CH,CH=C(CH,), 1.44k0.30 0 5.17+ 1.00 78.7k8.4 406-456 * CH,CH=C(CH,), 0.24f0.32 0 5.174 1.00 78.7k8.4 406-456 cyclopentene -0.80k0.15 - 10.6k 1.3 1.40k0.77 61.1 k7.1 406-489 The rate parameters of reaction ( 5 ) are expressed per reaction site; an asterisk indicates the site of addition in an unsymmetrical olefin.Evaluated values of A_, and A, imply a fairly ‘loose’ transition state in reaction (5). The Si-C bond energy has been estimated. *SiCl, radicals have been revealed to be electrophilic and susceptible to steric hindrance.In Part 2‘ of this series we reported a kinetic study of trichlorosilyl radical additions to 1-olefins. Using the same techniques we have carried out a similar study on cis- bu t-2-ene, zrans- but-2-ene, cis-pen t -2-ene, 2-methyl-bu t-2-ene and cyclopentene. In addition to the rate parameters, the orientation of addition in unsymmetrical olefins has been examined carefully because it may give some information about the nature of the transition state.2 11411142 ADDITION OF TRICHLOROSILYL RADICALS EXPERIMENTAL Trichlorosilane and acetone were purified as described previ~usly.~ Reagent grade cis- but-2-ene, trans-but-2-ene (Takachiho Kagaku) were used as received. cis-Pent-2-ene, 2- methyl-but-2-ene and cyclopentene were vacuum distilled. The purities of these olefins were checked by g.1.c.The sample preparations, photolyses and analyses were performed on a conventional greaseless and mercury-free vacuum line. The light from a 50 W medium-pressure mercury arc was filtered (Toshiba UV-33) and admitted to a cylindrical quartz reaction vessel of 139 cm3. Once photolysed, the reaction mixture was allowed to diffuse to a gas-sampling loop and was immediately analysed by a gas-chromatograph with a Gow-Mac gas-density detector. The adduct from each olefin was identified mainly by g.c.-m.s. analyses, which also made it possible to distinguish the isomeric adducts from an unsymmetrical olefin. In addition, the g.1.c. retention time of an adduct was compared with that of the authentic sample prepared by a radiation-induced reaction of trichlorosilane with the respective 0lefi1-1.~ RESULTS AND DISCUSSION Trichlorosilane, acetone and one of the 2-olefins (cis-but-2-ene, trans-but-2-ene, cis-pent-2-ene, 2-methyl-but-2-ene and cyclopentene) were mixed and photolysed at ca.440 K. Reaction products were almost exclusively the adducts of trichlorosilane to acetone and the respective olefin. The adduct to either end of the double bond was produced when an unsymmetrical olefin was used. The photoirradiation time was arranged to keep the reactant consumption < 5%. To suppress the formation of the higher telomers,l the ratio of olefin to trichlorosilane was kept as small as possible. Photolysis of acetone in the presence of HSiCl, and olefin has been shown to proceed via the following free-radical chain reactions3 hv CH,COCH, + 2 CH3 + CO CH, + HSiC1, (TCS) + CH, + SiCl, SiCl, + CH,COCH, (Ac) + (CH,)$OSiCl, \ / \ I SiCI, + C=C, (01) s ,C-<i-sicl, / (CH,),~OSiCl, + HSiCl, + (CH,),CHOSiCl, (A) + SiCl, \ I I I $-~-SiC13 + HSiCl, + HF-7-SiCl, (0) + SiCI,.In the scheme above, two chain cycles occur concurrently and result in the adducts A and 0. Assuming low conversions of the reactants and long reaction chains, a conventional steady-state treatment of the above scheme leads to the following rate equation (7) (RA/RO) ([olI/[Acl) = (k3/k5) {' + (k-5/k6) [TCS1-'l where k, denotes the rate constant of reaction (x). The rates of formation of A(R,) and 0 (R,) were measured as a function of TCS concentrations at various temperatures. They are given in table 1 in the form convenient for testing the validity of eqn (7).Acetone concentrations were (0.27-2.7) x mol cm-, and the ratio of olefin to acetone was from 0.1 to 0.3. From cis-pent-2-ene both 2-trichlorosilylpentane and 3-trichlorosilylpentane were formed. The ratio of the rate of formation of the two isomers was extensively studied and was found to be &/R3 = 0.471 k0.033. The ratio was virtually constant underT. DOHMARU AND Y. NAGATA 1143 TABLE 1 .-DEPENDENCE OF (R,/R,) ([Ol]/[Ac]) (b) ON [HSiCl,]-' (a) FOR trans-sUT-2-E~E, TEMPERATURES (U = 1 O5 cm3 mol-l) CiS-BUT-2-ENE, CiS-PENT-2-ENE, 2-METHYL-BUT-2-ENE AND CY CLOPENTENE AT VARIOUS 0 \*- / -\ \- \=/ \ * / \ -\ T/K alU b T / K a/U b T / K a / U b T / K a / U b T / K a/U b ~~ 378 378 378 378 404 404 404 404 416 416 416 416 423 423 423 423 428 428 428 428 433 433 433 433 438 438 438 438 443 443 443 443 448 448 448 448 452 452 452 452 477 477 477 477 50 1 50 1 50 1 50 1 50 1 ~~ 1.86 0.075 3.72 0.085 5.31 0.083 7.44 0.082 1.86 0.091 3.72 0.132 5.31 0.140 7.44 0.169 1.82 0.106 3.64 0.154 5.21 0.190 7.28 0.229 1.82 0.125 3.64 0.184 5.21 0.228 7.28 0.276 1.86 0.145 3.72 0.221 5.31 0.276 7.44 0.356 1.22 0.141 3.64 0.250 5.21 0.331 7.28 0.412 1.82 0.177 3.64 0.300 5.21 0.419 7.28 0.532 1.82 0.195 3.64 0.387 5.21 0.501 7.28 0.680 1.82 0.233 3.64 0.474 5.21 0.608 7.28 0.820 1.86 0.306 3.72 0.545 5.31 0.689 7.44 1.036 1.86 0.760 3.72 1.43 5.31 1.99 7.44 2.85 1.86 1.52 2.48 1.99 3.72 2.90 5.31 4.27 7.44 5.98 397 1.77 0.075 397 3.48 0.091 397 7.04 0.105 418 1.77 0.0953 418 3.53 0.123 418 5.29 0.146 418 7.07 0.194 443 1.77 0.178 443 3.55 0.266 443 5.32 0.379 443 7.09 0.495 452 1.77 0.235 452 3.53 0.403 452 5.32 0.579 452 7.09 0.754 462 1.77 0.326 462 3.55 0.577 462 4.16 0.663 462 5.32 0.810 462 5.92 0.959 462 7.09 1.129 406 1.76 0.185 406 3.51 0.231 406 5.01 0.289 406 7.02 0.342 423 1.76 0.281 423 3.51 0.419 423 5.01 0.554 423 7.02 0.728 439 1.76 0.464 439 3.51 0.881 439 5.01 1.168 439 7.02 1.449 455 1.76 0.773 455 3.51 1.671 455 5.01 2.373 455 7.02 3.128 406 1.89 0.089 406 1.76 0.148 406 3.78 0.167 406 3.51 0.144 406 5.40 0.206 406 5.01 0.166 406 7.56 0.363 406 7.02 0.164 423 1.89 0.203 423 1.76 0.155 423 3.78 0.366 423 3.51 0.198 423 5.40 0.476 423 5.01 0.198 423 7.56 0.650 423 7.02 0.192 440 1.89 0.465 439 1.76 0.179 440 3.78 0.902 439 3.51 0.232 440 5.40 1.34 439 5.01 0.285 440 7.56 1.88 439 7.02 0.294 448 1.26 0.423 456 1.76 0.265 448 1.51 0.579 456 3.51 0.319 448 1.89 0.746 456 5.01 0.361 448 2.52 0.948 456 7.02 0.474 448 3.78 1.42 472 1.76 0.323 448 4.45 1.79 472 3.51 0.523 448 5.40 2.11 472 5.01 0.653 448 7.56 3.08 472 7.02 0.803 456 1.89 1.13 489 1.76 0.550 456 3.78 2.21 489 3.51 0.895 456 5.40 3.02 489 5.01 1.15 456 7.56 4.06 489 7.02 1.51 * Indicates the site of addition.An adduct to the other site is not shown here; see text.1144 ADDITION OF TRICHLOROSILYL RADICALS the present experimental conditions. The values of b for the 2-isomer are easily calculable from those for the 3-isomer using RJR, and are omitted from table 1. Speir and Webster, studied peroxide-initiated reactions of trichlorosilane with pent-2-ene between 353 and 368 K in the liquid phase and suggested the formation of the two isomers with R,/R, = 2.3.It is of interest that the value in the gas phase is much different from that in liquid phase, but the difference cannot be explained at the present time. A detailed study of the dependence of the ratio on phase and temperature is continuing in our laboratory. 2-Methyl-but-2-ene gives 94.1 _+ 1.6% 2-methyl-3-trichlorosilylbutane and 5.9 f 1.6% 2-methyl-2-trichlorosilylbutane; the ratio of the two isomers is also invariant under these experimental conditions. No experimental value is available for this ratio. Values (RJR,) ([Ol]/[Ac]) and ITCS1-l in table 1 were plotted according to eqn (7); the linearity of the plots at each temperature was excellent for all the olefins studied.Eqn (7) shows that the intercept of the linear plot gives k,/k, and the slope gives (k3/k5) ( k 5 / k 6 ) . Accordingly, k-,/k, is obtained from the slope divided by its intercept. Arrhenius parameters of k,/k, and k-,/k, for each substrate were calculated by a weighted linear least-squares method.,* The results are given in tabular form in the abstract; error limits are the standard deviations from least-mean-squares plots. EVALUATION OF ELEMENTARY RATE PARAMETERS Hydrogen abstraction reactions by alkyl radicals have been extensively studied so that A factors of this type of reaction are generalized to be 1011.5*0.5 cm3 mol-l ~ - 1 . ~ Recently, Arthur et aL8 evaluated the kinetic data for hydrogen abstraction from silanes.From their compilation we remove the A factor for hydrogen abstraction from HSiCl, and obtain as an average value 1011.6*o.8 cm3 mol-l s-l. Therefore it is reasonable to assume that the A factor for reaction (6) is The attacking radicals in reaction (6) are p-substituted alkyl radicals, and may be considered to have small polarities, since the substituent effect is attenuated by the -CH2- groupO9 In this narrow series of very similar reactions, the Polanyi relationlo may be used to estimate the activation energies. Activation energies for the reaction of -CH, (18 kJ mo1-l)l1 and -CH,CH, (22.2 kJ mol-l)ll with HSiCl, were adopted as standards. The enthalpy changes were calculated from D(H-SiCl,) = 382 kJ mol-l obtained by Walsh12 and the well-established values of D(R-H).13 Using the k , values thus obtained, k-, values may be calculated as shown in table 2.The values for 1-olefins previously obtained, are also presented in table 2 for the purpose of the following discussion. The values of log A _ , for 2-olefins other than cyclopentene are much greater than 13.4, the value giving ASS = 0 at 440 K. It follows that SiCl, detachment and consequently SiCl, addition proceed through fairly ' loose' transition states in the case of 2-olefins as well as 1-olefins. Cyclopentene behaves in a completely different way. More work on cyclic olefins would be needed to undertake a detailed discussion. (8) cm3 mol-l s-l. A , is related to A _ , by eqn (8): log (AJA-,) = AS5O/19.15 +log,, (e x RT) where AS: is a standard entropy change accompanying reaction ( 5 ) and log,, (e x RT) is a conversion factor for the change in standard state from 1 atm at s.t.p. to 1 mol cm-,.Entropies of the respective adduct radical may be calculated by the * In runs for the 2-olefins other than cyclopentene, the error limits of the intercept values were much larger than those of the slope values at higher temperatures. Thus, the intercept values were averaged first and the k-,/k, ratios were calculated using the averaged values.T. DOHMARU AND Y. NAGATA 1145 TABLE 2.-EVALUATED VALUES OF ELEMENTARY RATE PARAMETERS 15.0- 117.2 kJ mol-I/RT In 10 14.8- 108.4 kJ mol-l/RT In l o RTln 10 15.3- 113.0 kJ mol-l/RTIn 10 14.8- 107.1 kJ mol-l/RT In 10 16.3- 116.3 kJ mol-l/RT In 10 15.6- 102.1 kJ mol-l/RT In 10 15.5- 101.3 kJ mol-'/RTln 10 16.9- 110.9 kJ mol-l/RT In 10 16.9- 110.9 kJ mol-l/RT In 10 16.7- 105.9 kJ mol-l/RT In 10 16.7- 103.8 kJ rnol-l/RT In 10 12.9-86.2 kJ mol-I/RT In 10 ~ 12.7 12.3 12.7 12.1 13.7 13.2 13.1 14.8 14.8 14.0 13.9 10.1 a Revised value.SiCl, additions to ethylene have been reinvestigated recently14 because the Taken Statistically adjusted for the previous study3 may have been inaccurate due to small amount of a certain impurity. from ref (1). number of identical sites in the olefins. The starred site shows the reaction centre. method described by O'Neal and Benson.15 As a first step, So[CH,CH2CH(CH,)SiCl,] is calculated as 464 J K-l moI-l employing the bond-additivity relations for silicon compounds.ls* Then corrections are made to obtain So[CH,cHCH(CH,)SiCl,] : the electronic degeneracy was included as ( + R In 2) and the variation in the barriers of the two hindered rotations was taken as 2.1 x 2 J K-l mol-l (a two-thirds reduction in barrier height was assumed). Mass differences, changes in vibration and moment- of-inertia differences were negligibly small and there was no symmetry change.? It follOws that SO(CH,tHCHCH,) = 474 (171) J K-l mol-l.I SiCl, The value in parenthesis is C:, obtained in the same way. Similarly, So(CH,CH2cHCHCH,) = 5 14 (1 94) J K-l mol-1 I SiCl, SO(CH,CH,CHtHCH,) = 5 13 (194) J K-l mol-I I SiC1, * S0[CH3CH,CH(CH3)SiCl3] = 9 x So(C-H) + 3 x So(C-C) + S"(C-Si) + 3 x So(Si-C1) - 3R In 3 t The radical centre was assumed to retain its planar configuration, so that there is no change in the (symmetry) + R In 2 (optical isomer).number of optical isomers.1146 ADDITION OF TRICHLOROSILYL RADICALS and SO (4. -) = 459 (159) J ~ - 1 mol-1 SiCl, were obtained. As the entropies and heat capacities of eSiC1, and the olefins are available in the literature,' we can get AS: at 500 K for each reaction site, and accordingly determine the values of A, using eqn (8). The A, values are also listed in table 2, together with the values previously obtained for 1-olefins. A, can now be calculated from A,. The average value of log(A3/cm3 mob1 s-l) obtained from all the olefins in table 2 is 12.6f 1.0. It is of interest to compare this value with log A, (isobutene) since acetone and isobutene have similar structures.The difference between the two A factors is small but tends to imply that the transition state of isobutene is a little looser than that of acetone. si-c B O N D ENERGY Combining our experimental activation energies with the z-bond energies7 of olefins, the bond energy of Si-C may be obtained as follows: D(Si-C) = g - A z (9) where is the n-bond energy of an individual olefin defined as in eqn (10) : q(CRlR2=CR3R4) = D(H--CR,R,CHR,R,) - D(H-CRlR2CR,R4) = Afl(CR,R,CHR,R,) + Aq(CHR,R,tR,R,) - Afl(CR1R2=CR3R4) - Aq(CHRlR2CHR3R4). (10) The heats of formation in eqn (10) are found in the literature7$ l7 or easily calculated using the group additivity rule.8 The absolute value of E3 is not known at present, but we may assume tentatively that E3 = 12.6 kJ mol-l (3.0 kcal mol-l).Then E, is provided from the table in the abstract. A% is given as E, - E-, + AnRT where AnRT is a conversion factor for the change in standard state; in this case An = - 1. The value of D(Si-C) thus obtained may be classified according to the number of substituents at the carbon to which the trichlorosilyl group attaches: primary, secondary or tertiary. Data from 1-olefins give D(Si-Cpri,) = 352+5 kJ mol-', data from (CH,),C=CHCH, give D(Si-C,,,,) = 325 kJ mol-1 and data from the rest of the 2-olefins give D(Si-C,,,) = 332 & 4 kJ mol-l. D(Si-CPri,) is in good agreement with D(Si-C) = 356f 17 kJ mol-l obtained by Potzinger et aZ.18 from electron-impact experiments combined with thermochemical calculations. D(Si-C,,,) and D(Si-C,,,,) seem a little too small compared with D(Si-Cpri,).All these values may be the upper limit of our data because E3 = 12.6 kJ mol-l, which this bond-energy calculation is based upon, may probably be the lower limit. * REACTIVITY OF *SiC13 RADICALS The rate constants for the addition of SiCl, radicals to olefins are compared with those of some other radicals in table 3. The O(3P) atoms and *NF2 radicals displayT. DOHMARU AND Y. NAGATA 1147 TABLE 3.-&LATIVE RATE CONSTANTSa OF ADDITION OF RADICALS TO OLEFINS - SiC13b 9 NFZC 0 (3P)d CH3e substrate (460 K) (373 K) (298 K) (453 K) CH,COCH, = f =/ -/\ - -/\/ - -/ -\ \=/ \- -\ \=/\ \-/ -\ Q 0.3 1 3.1 3.4 4.1 13.4 4.7 3.7 6.5 16.7 1.4 1 4.4 4.3 19.6 10.3 10.8 33.6 7.9 1 5.8 5.8 1 0.7 0.7 25 1.1 24 28 0.2 0.4 79 0.4 27 a Rate constants per double bond (not per reaction site) were employed for comparison.This work and ref (1). A. J. Dijstra, J. A. Kerr and A. F. Trotman-Dickenson, J. Chem. Soc. R. J. Cvetanovic and A, 1967, 105. R. J. Cvetanovic, Adu. Photochem., 1963, 1, 161. R. S. Irwin, J. Chem. Phys., 1967, 46, 1694. f Reinvestigated value. typical electrophilic trends and the CH, radicals display much less se1e~tivity.l~ SiCl, radicals behave similarly to O(,P) and *NF, for all the 1-olefins. In view of the trend in the 2-olefins, however, the increment in reactivity due to one alkyl group substitution becomes 2 or 3 times smaller than in the 1-olefins for the case of *SiCl, radicals. No such trend is observed in the case of O(,P) atoms. These trends may be interpreted in terms of the steric effect of the bulky *SiCl, radical.The orientations of addition to the unsymmetrical 2-olefins were not affected on changing the reaction temperature from 406 to 456 K; however, this fact alone does not provide any decisive information on the configuration of the transition state. T. Dohmaru and Y. Nagata, J. Chem. SOC., Faraday Trans. 1, 1979, 75, 2617. T. Dohmaru, Y. Nagata and J. Tsurugi, Chem. Lett., 1973, 1031. A. M. El-Abbady and L. C. Anderson, J. Am. Chem. SOC., 1958,80, 1737. J. L. Speier and J. A. Webster, J. Org. Chem., 1956, 21, 1044. R. J. Cvetanovic, R. P. Overend and G. Paraskevopoulos, Znr. J. Chem. Kinet., 1975, S1, 249. S. W. Benson, Thermochemical Kinetics (Wiley, New York, 2nd edn, 1976). N. L. Arthur and T. N. Bell, Rev. Chem. Intermediates, 1978, 2, 37. S. H. Marcus, W. F. Reynolds and S. I. Miller, J. Org. Chem., 1966, 31, 1872. lo J. A. Kerr, in Free Radicals, ed. J. K. Kochi (Wiley, New York, 1973), chap. 1. l1 J. A. Kerr, A. Stephens and J. C. Young, Znt. J. Chem. Kinet., 1969, 1, 371. l 2 R. Walsh and M. J. Wells, J. Chem. Soc., Faraday Trans. I , 1976, 72, 1212. l3 G. A. Russel, in Free Radicals, ed. J. K. Kochi, (Wiley, New York, 1973), chap. 7. l4 T. Dohmaru and Y. Nagata, to be published. l5 H. E. O’Neal and S. W. Benson, in Free Radicals, ed. J. K. Kochi (Wiley, New York, 1973), chap. * J. M. Tedder and J. C. Walton, Acc. Chem. Res., 1976, 9, 183. 17. H. E. O’Neal and M. A. Ring, Znorg. Chem., 1966, 5, 435.1148 ADDITION OF TRICHLOROSILYL RADICALS *' S. W. Benson, F. R. Cruickshank, D. M. Golden, G. R. Hauger, H. E. O'Neal, A. S. Rodgers, R. Shaw and R. Walsh, Chem. Rev., 1969, 69, 279. P. Potzinger, A. Ritter and J. K. Krause, 2. Naturforsch., Teil A , 1975, 30, 347. (Butterworths, London, 1972). 19 J. A. Kerr and M. J. Parsonage, Eualuated Kinetic Data on Gas Phase Addition Reactions (PAPER 1 /753)

ISSN:0300-9599    

DOI:10.1039/F19827801141

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1, 1982, 78, 1149-1 164 Kinetics of Metal Oxide Dissolution Reductive Dissolution of Nickel Ferrite by Tris(picolinato)vanadium(rI) BY MICHAEL G. SEGAL AND ROBIN M. SELLERS* Central Electricity Generating Board, Berkeley Nuclear Laboratories, Berkeley, Gloucestershire GL 13 9PB Received 1 lth May, 1981 A detailed study of the reductive dissolution of NiFe,O, by V@ic); is reported. The kinetics of the reaction follow a cubic rate law and exhibit a first-order dependence on [V'I], indicating that the rate-determining step involves attack by V@ic); on Fe"' ions at the oxide surface. Dependences on [H+] and [free picolinate] are attributed to adsorption of these species at surface sites according to a simple model based on the Langmuir adsorption isotherm.Buffer and surfactant concentration had little effect. Some pitting etc. of the oxide surface occurs as the dissolution proceeds, although there appears also to be some general surface attack. The nature of the surface sites and subsequent steps in the dissolution process are discussed. The dissolution of metal oxides is of considerable practical importance in fields as diverse as the extraction of metals from ores, the transport of metals in the hydrologic cycle and the removal of oxide deposits from metal surfaces. Many metal oxides are difficult to dissolve, even in concentrated acids or in the presence of strong complexing agents. Our interest in this problem stems from the relatively poor oxide removal achieved in cleaning the pipework of water-cooled nuclear reactors, and in particular the removal of nickel ferrite, NiFe,O, (trevorite), and similar materials from the primary circuit of pressurised-water reactors.Nickel ferrite is often considered to be insoluble in acid,l although dissolution can be effected after heating in, for instance, phosphoric Complexing agents such as EDTA do not appear to have any significant effect on the rate of such reactions., In this paper we focus attention on the use of reducing agents to bring about dissolution, and describe a detailed study of the kinetics of dissolution of stoichiometric nickel ferrite by tris(pico1inato)vanad- ium(I1) in aqueous solution. Little previous work on reductive dissolution has been published. The method is perhaps most familiar in the dissolution of MnO, by acidified hydrogen peroxide, in the dissolution of anhydrous CrCl, by CrC1,,4 and in the use of thioglycollic acid in the analysis of iron(1n) oxide~.~ It has also been used to characterise soil samples by extraction of Fe and A1 using dithionite and oxalate,6 and patents have been issued for the dissolution of iron oxide scales using reductants such as erythorbic acid,' ascorbic acid7 and thioglycollic acid.8 More detailed studies have been made by Zabin and Taubeg on the reaction of Cr2+ with oxides such as MnO,, PbO,, Tl,O,, Mn,O,, Co,O, and CeO, by ValverdelO on the dissolution of FeO, Fe,O,, Fe,O,, COO, CuO, etc.by various metal-ion redox couples, by Bradburyll on the dissolution of Fe,O, by thioglycollic acid, etc. and our own preliminary work on the dissolution of FelI1 oxides by complexes of CrII, VII, FeII, etc.12 Of related interest is the work of Pryor and Evans1, and others14 on the electrochemical reduction of oxides on metal electrodes.11491150 METAL OXIDE DISSOLUTION KINETICS EXPERIMENTAL MATERIALS All solutions were prepared from triply distilled water. Picolinic acid (Aldrich) and sodium formate (B.D.H.) were recrystallised once from water before use. V2+ was prepared by electrolytic reduction of 0.2 mol dm-3 VOSO, in 0.5 mol dm-3 H,SO, at a lead cathode and stored under argon. The total vanadium content of the V" stock solution was determined spectrophotometrically as the peroxovanadium(v) complex taking = 28 1 dm3 mol-l cm-l, and the amount of V2+ and V3+ by direct absorbance measurements on the stock solution at 400 nm, taking&(V3+) = 8.3 dm3 mol-l cm-l and &(V2+) = 0.9 dm3 mol-l cm-l.The surfactants Triton X-100 and Hyamine 1622 were B.D.H. G.P.R. grade. All other chemicals were AnalaR grade and used without further purification. Nickel ferrite was prepared either by precipitation of the oxalates from a mixture containing the appropriate quantities of nickel and ferrous acetate (the 'oxalate' method)', l5 or by the precipitation of the hydroxides by addition of Na,C03 to a mixture of nickel and ferric nitrates (the 'carbonate' method). Precipitates were washed with triply distilled water (ca. 15 dm3 per 15 g of sample) to remove all acetic acid+NaNO,, dried in an oven at 110 OC for 1-2 d, and finally calcined at 1000-1400 OC under argon to produce the ferrite. Sizing of the oxide particles was achieved by grinding and sieving using Endicott sieves.The purity of the oxides was checked by infrared spectroscopy (Perkin Elmer 180 photometer) and X-ray diffraction (Philips 2 kW X-ray generator). No evidence for impurities such as NiO or Fe203 was found, except in samples C1 and C2 (see table 1) which were calcined below the ferrite transition temperature. The lattice parameters measured agree well with literature values.16 Chemical compositions were confirmed by dissolving the oxides in phosphoric acid (20 mg oxide in 50 cm3 B.D.H. AristaR H3PO4, refluxed for 12 h) or hydrochloric acid (20 mg oxide in 50 cm3 B.D.H. AristaR HC1, refluxed for 24 h), and determining the iron and nickel contents by atomic absorption.Analyses found were sample A1 (in H3P0,):44.1% Fe, 25.4% Ni; A1 (in HC1):44.3% Fe, 22.2% Ni; B6 (in H3P0,):48.5% Fe, 27.4% Ni; C3 (in H3P0,):48.2% Fe, 24.5% Ni; calculated for NiFe,0,:47.7% Fe, 25.0% Ni. Similar results were obtained in the dissolution of these samples by V(pic);.* APPARATUS AND PROCEDURE The kinetic runs were carried out in glass reaction vessels. Constant temperature (in the range 40-80 "C) was maintained by circulating water from a water bath through a jacket around the reaction vessel. Solutions of picolinic acid, etc. containing ca. 2 mg of oxide were degassed by bubbling with high-purity argon (B.O.C.) for i-1 h before initiating the reaction by addition of V2+. The !!(pic); reagent is sensitive to oxygen, and to ensure no complications from aerial oxidation a slow stream of argon was passed through the solution during the run.Solutions were stirred using a magnetic stirring bar. The rate of dissolution was measured by withdrawing CQ. 1.5 cm3 aliquots by syringe at regular intervals. The dissolution reaction was arrested by exposure to air, the V(pi~)~ produced being unreactive towards nickel ferrite (see below). The samples were then diluted x 10 with 0.1 mol dm-3 HC1 (B.D.H. C.V.S. reagent), filtered through a 0.22 pm filter (Millipore Millex GS) to remove any unreacted oxide, and analysed for iron and nickel content by atomic absorption spectrometry using a Baird A5100 instrument. At the end of each run some of the reakent was cooled to room temperature under argon, and the pH measured on an EIL type 7050 meter.Surface area measurements were made by N, absorption with a Perkin Elmer 212D sorptometer by the B.E.T. single-point method. Scanning electron micrographs were taken with om; * pic- = picolinateM. G. SEGAL A N D R. M. SELLERS 1151 a Cambridge Stereoscan 150. Partial dissolution experiments to investigate changes in surface morphology were done using ca. 25 mg oxide and appropriate reagent concentrations. After the required amount of dissolution had taken place the oxide was removed by filtration, washed with 20 cm3 HC1 and 50 cm3 H,O and dried in an oven at 110 OC for ca. 1 h before mounting for observation in the scanning electron microscope. RESULTS MORPHOLOGY OF THE PARTICLES The morphology of the nickel ferrite particles was characterised by surface area determinations and scanning electron microscopy.The surface area measurements are summarised in table 1. For material calcined at 1400 O C the figures are only slightly TABLE 1 .-CALCINING HISTORIES AND CHARACTERISTICS OF NICKEL FERRITE SAMPLES specific lattice sample preparative calcining particle surface area parameter no.a method conditions size/pmb /m2 g-lc %/Ad A1 A2 A3 Bl B2 B3 B4 B5 B6 c1 c2 c 3 c 4 c 5 C6 oxalate oxalate oxalate carbonate carbonate carbonate carbonate carbonate carbonate carbonate carbonate carbonate carbonate carbonate carbonate 6 h at 1000 OC 12 h at 1000 O C 24 h at 1400 OC 6 h at 1400 O C 6 h at 1400 OC 6 h at 1400 OC 6 h at 1400 OC 6 h at 1400 OC 6 h a t 1400OC 6 h at 600 OCe 6 h at 800 O C e 6 h at 1000 OC 6 h at 1200 O C 6 h a t 1400OC 72 h at 1400 OC < 100 < 100 < 100 > 300 150-300 106-150 75-106 53-75 < 53 < 53 < 53 < 53 < 53 < 53 < 53 1.2 - < 0.04 ca.0.08 0.14 8.3374 - - - 8.3383 - 8.3395 - a A2 and A3 prepared by recalcining Al; B1-B6 calcined as one sample, and then sieved; by from X-ray powder C1C6 same batch divided after drying at 110 OC, and calcined and sieved separately; sieving, sizes correspond to particle diameters; diffraction; by N, adsorption; X-ray and infrared measurements show these samples to contain some Fe,O,. larger than the calculated geometric value, which for spherical particles with a radius of 10 pm is 0.056 m2 g-l taking the density as 5.37 g Lowering the calcining temperature increased the specific surface area by a factor of ca.10, indicating some porosity in the oxide formed under these conditions. This conclusion was borne out by the scanning electron microscopy results (plates 1 and 2) which showed the 1400 O C calcined material to consist of amorphous particles, with more or less flat faces. Under high resolution stepped structures were detectable (plate 2). It is also noteworthy that these larger particles (typically > 50 pm in diameter) have associated with them some much smaller particles (< 1 pm in diameter), which have not been separated by sieving. The particles of material prepared by the oxalate method and calcined at1152 METAL OXIDE DISSOLUTION KINETICS 1000 O C were to good approximation spherical, but with a much more convoluted surface morphology.The nature of the surface changes during dissolution. At least three types of attack can be recognised, although not every particle exhibits the features of all three. In the first regular pits are formed as illustrated in plate 3(a). In cross-section these are hexagonal, triangular, rectangular or octagonal, with hexagonal pits being perhaps the most common. Deep fissures are formed in other particles, probably as a result of attack at grain boundaries or edge dislocations. This behaviour is shown in plate 3(b). Finally there is a more general surface attack, apparently not directed at any specific site. Plate 4(a) illustrates attack of this type on what seems initially to have been a stepped structure. One or two particles also showed a ‘dendritic’ like attack as shown in plate 4(b).This probably results from the dissolution of an iron-rich phase, and is indicative of some microinhomogeneity in a small fraction of the particles. VARIATION OF THE AMOUNT OF OXIDE DISSOLVED WITH TIME The variation of the amount of nickel ferrite dissolved with time in a typical dissolution run is shown in fig. 1. The main features are a rapid initial dissolution rate, I I 1 I I 1 100 150 200 250 0 50 time/min FIG. 1 .-Variation of amount of oxide dissolved with time. Measurements made in solutions containing 3.2 x mol dm3 V(pic); + 0.030 mol dm-3 free picolinate + 0.1 mol dm-3 HCO; + 20 ppm Triton X-100, pH 4.3, T = 80 OC and using nickel ferrite sample Al. 0, percentage dissolution in terms of the amount of Fe dissolved; 0, based on Ni dissolved.Line calculated according to eqn (7) with koba = 4.5 x min-l. falling off as the size of the particles decreases. This behaviour is interpreted on the assumption that reaction occurs at the surface of the particles, at a rate proportional to the instantaneous surface area, i.e. dM dt - = - k A where M is the mass of undissolved oxide, k is a rate constant in g m2 min-l and A is the total surface area of the oxide in m2. For spherical particles of uniform sizeJ . Chem. Soc., Faraday Trans. 1, Vol. 78, part 4 Plate 1 PLATE 1 .-Scanning electron micrographs of NiFe,O, samples calcined at lo00 OC. (a) Sample A1 ; prepared by the oxalate method. (b) Sample C3; prepared by the carbonate method. M. G. SEGAL AND R.M. SELLERS (Facing p . 1 152)J. Chem. SOC., Faraday Trans. 1, Vol. 78, part 4 Plate 2 PLATE 2.-Scanning electron micrographs of NiFe,O, samples calcined at 1400 O C . (a) Sample B5; (b) sample B4, detail showing stepped structure. M. G. SEGAL AND R. M. SELLERSJ . Chem. Soc., Faraday Trans. 1, Vol. 78, part 4 Plate 3 PLATE 3 .-Scanning electron micrographs of NiFe,O, sample B5 after treatment in solution containing 1.5 x lop2 mol dmp3 V(pic);+3.2 x lo-* mol dmP3 free picolinate+0.2 mol dmP3 HCO;+ 100 ppm Triton X-100 at 80 OC for 5 h (25% dissolved). (a) Regular pits; (6) grain boundary attack. M. G. SEGAL AND R. M. SELLERSJ. Chem. SOC., Faraday Trans. 1 , Vol. 78, part 4 Plate 4 PLATE 4.-Scanning electron micrographs of NiFe,O, particles after treatment in solution containing 1.5 x lo-, mol dm-3 V(pic); + 3.2 x lo-, mol dmP3 free picolinate+0.2 mol dmP3 HCO;+ 100 ppm Triton X-100 at 80 OC.(a) Sample B5 after 35 min in reagent (8.4% dissolved) showing result of general surface attack on stepped structure; (b) sample B5 after 5 h in reagent (25% dissolved) showing dendritic attack. M. G. SEGAL AND R. M. SELLERSM. G. SEGAL A N D R. M. SELLERS 1153 (radius, r), the surface area, A , will be given by eqn (2), where N is the number of particles, and the total mass, M, is given by eqn (3) A = 47rr2N (2) (3) 4 3 M = -nr3 Np. Substituting eqn (2) and (3) into eqn (1) and integrating gives eqn (4), where Mo is the mass of the particles at t = 0 Experimentally it is convenient to measure the concentration, Ct, of the metal ion (Fe or Ni) in solution, which is given by eqn ( 5 ) where x is the weight fraction of the metal in the oxide and V is the volume of the solution.When all the oxide has dissolved eqn (5) becomes eqn (6) XMO C , =- v ' Substituting eqn (5) and (6) into eqn (4) and rearranging gives eqn (7) ( l - z y = 1 - & kt (7) where ro is the initial particle radius. Thus a plot of [ 1 - (C,/C,)]i against t should give a straight line with a slope equal to kobs (= k/rop), an intercept of 1 on the [ 1 - (Ct/C,)]i axis and an intercept of k& (= tm) on the time axis. Fig. 2(a) shows the data of fig. 1 treated in this way, from I I I I I I 0 0.2 0.4 0.6 0.8 1.0 I I 1 I I I 0 0.2 0.4 0.6 0.8 1.0 tlt, FIG. 2.--Cubic rate law plots. (a) 0, Data of fig.1 replotted according to cubic rate law, eqn (7), t , = 220 min; 0, cubic rate law plot for dissolution of nickel ferrite sample with narrow particle size distribu- tion (sample B3), reagents and conditions as for fig. 1, t , = 1 1 0 0 min. (b) Calculated cubic rate law plot for a non-uniform size distribution (see text for details). 38 FAR 11154 METAL OXIDE DISSOLUTION KINETICS which it is seen that a linear dependence according to eqn (7) is obtained up to ca. 75% dissolution {[ 1 - (C,/C,)]i = 0.63). Thereafter the dissolution becomes slower than predicted by this equation, because the particles do not in reality have a uniform size. That this is so we have been able to demonstrate in two ways. The first is by measuring the dissolution rate as a function of time for an oxide sample with a very narrow size distribution (sample B3, table 1).This obeyed the cubic rate law to at least 90% dissolution, as shown in fig. 2(a). Secondly we have calculated the variation of [ 1 - (Ct/Cw)]i with time on the assumption that the oxide consists initially of equal masses of particles of radius r, 2r, 3r, 4r and 5r. As shown in fig. 2(b) this is non-linear over long times, but approximates to a straight line for the first ca. 60% of the dissolution {[l -(Ct/Cw)]i 2 0.731, and for practical purposes behaves as if the distribution were uniform, with radius ca. 2.5- [straight line, fig. 2(b)]. The validity of eqn (7) has also been investigated by determining kobs as a function of initial particle size. For this purpose a sample of nickel ferrite was sieved into six size fractions (cf.table l), and the dissolution kinetics measured under otherwise constant conditions. The results of this experiment are shown in fig. 3, from which it 15 I c ... ,E 10 Ol I 1 I 1 J 0 100 200 300 400 500 particle diameterlpm FIG. 3.-Effect of particle size on the dissolution kinetics of NiFe,O, by V(pic);. Reagents and conditions as for fig. 1, but using oxide samples B 1-B6. Vertical error bars are mean deviations, horizontal error bars cover range of particle sizes as determined by sieving. Line is calculated according to eqn (8) with k / p = 7 x lo-* m min-l. is clear that there is an inverse relationship between kobs and particle diameter, d. The line shown in fig. 3 is drawn according to eqn (8) with k / p = 7 x m min-l, and gives a reasonable fit to the data Plotting k& against d yields a straight line, but with intercept k& = ca, 300 min when d = 0, indicating some agglomeration of smaller particles.The cubic rate law, eqn (7), clearly provides a good description of the kinetics of th? V(pic); + NiFe,O, reaction, and establishes that the rate-determining step involves reaction at the particle surface. For almost all the work described here dissolution rates were determined from the amount of iron in solution. Rates based on nickel gave similar results (cf fig. 1 and 6). Detailed comparison of the two reveals, however, that the iron : nickel ratio varies from the stoichiometric value of 2 in the early stages of the dissolution (< 30%), and, as the data in fig.4 show, iron is dissolved preferentially.M. G. SEGAL AND R. M. SELLERS 1155 0 1 I I I I I I I 0 10 20 30 40 50 60 time/min FIG. 4.-Variation of the iron: nickel ratio in solution as a function of time in the dissolution of NiFe,O, by V(pic);. Measurements made in solutions containing 3.2 x mol dm-3 V(pic); + 3.0 x lo-, mol dm-3 free picolinate + 0.10 mol dmd3 HCO; + 20 ppm Triton X-100, pH 4.2, T = 80 O C and using nickel ferrite sample A1 . EFFECT OF REAGENT CONCENTRATION ON THE DISSOLUTION KINETICS Determination of the effect of reagent concentration on dissolution kinetics was carried out with oxide sample Al. The solutions contained in general an excess of picolinate to ensure that all V2+ was present as the blue-black coloured 1 : 3 complex, although in a few runs a significant amount of the deep red 1 :2 complex was present.The concentration of uncomplexed picolinate in the solution was calculated on the basis of complete formation of the 1:3 complex, or taken to be zero if the total [picolinate] was less than 3[V2+]. The stability constants for the formation of these complexes are log Kl = 4.4, log K2 = 4.6 and log K3 = 3.8 in 0.5 mol dm-3 KCl at 25 'C.la Corrections were also made to take into account the V3+ present as an I I 1 I 1 I 1 1 0 2 4 6 8 10 12 14 [ W]/ 1 0-3 mol dm-3 FIG. 5.-Effect of vanadous concentration on the kinetics of-dissolution of NiFe,O, by V(pic);. 0, Measurements made in solutions containing 0.020 mol dm-3 free picolinate + 0.1 mol dmw3 HCO;+20 ppm Triton X-100, pH 4.0, T = 80 OC and using nickel ferrite sample Al.0, Measurements made in solutions containing 0.01 8 rnol dm-3 free picolinate + 0.1 mol dm-3 CH3CO; + 20 ppm Triton X-100, pH 4.4, T = 0 OC and using nickel ferrite sample Al. Lines calculated according to eqn (1 l), using constants shown in table 7. 38-21156 METAL OXIDE DISSOLUTION KINETICS impurity (typically 3-779 in the V2+ solution, again assuming complete formation of the 1 : 3 complex, for which log Q3 = 1 5.4.lS Solutions also contained a surfactant to aid dispersion of the oxide and a buffer to keep the pH constant. The buffer concentration was 0.1 mol dm-3 in most runs and effectively maintained the ionic strength of the solution constant at ca. 0.1 mol dm-3. Reaction rate constants, kobs, were calculated from the slopes of plots of [l -(Ct/Cm)$ against time [ i e .according to the cubic rate law, eqn (7)]. Values of C, were based on the weight of oxide present initially. In a number of runs where the dissolution rate was rapid, the reaction was allowed to go to completion, and C , determined experimentally. Agreement with calculated values was within & 10%. The concentration dependences found are shown in fig. 5-7 and tables 2-4, and can be summarised as follows. VII Using formate as a buffer (80 "C) a good linear dependence on [V1] was found (fig. 5 ) for [VII] = (1 - 13) x mol dm-3, with no evidence for any pathway indepen- dent of [VII]. With acetate as the buffer (60 "C) a linear dependence on [VII] was again found (fig. 5) for [VII] = (0.9-7) x mol dm-3.Here, however, there was some indication of a [VI1]-independent pathway. A dependence of the form given in eqn (9) provides a good description of the data, and a linear regression of [VII] on kobs yields a = 3 x lo-* min-l and b = 0.4 dm3 mol-l mip-l kobs = a 4- b[V'I]. (9) However, the intercept falls within the relatively large experimental errors (ca. 15 %) involved in these experiments, and the measurements do not therefore unambiguously establish the existence of a [VIII-independent pathway. Other interpretations of the data are possible (for instance that the reaction becomes less than first order in [VII] at high [VII]), but we conclude that under the conditions of our experiments the only pathway of any consequence is that involving direct reaction of VII with the surface. PICOLINATE Only a weak dependence on [picolinate] was found, kobs falling by a factor of ca.2 on increasing the free picolinate concentration from 0 to 0.35 mol dm-3 in the presence of formate buffer (80 "C) or by a factor of ca. 6 in the concentration range 0-0.2 mol . dm-3 with acetate buffer (60 "C) as shown in fig. 6. PH A marked increase in rate was found on decreasing the pH. With formate buffer (80 "C) the rate increased ca. 20-fold in the pH range 6-3.2 (fig. 7) and ca. 10-fold in the same range in the presence of acetate buffer (60 "C); pH levels were adjusted by addition of HCl or NaOH, and to check that the counter ions (Cl-, Na+) had no effect some runs were done in which excess NaCl or Na2S0, was added. No significant change in rate was found (cf.table 4). BUFFER With formate as the buffer the rate was independent of [HCO,-] for [HCO,-] = 0.1 - 0.5 mo1.dm-3 (table 2). Substituting acetate or cacodylate as the buffer gave essentially the same results.M. G. SEGAL AND R. M. SELLERS 2.0k 1157 - FIG. 6.-Effect of picolinate concentration on the kinetics of dissolution of NiFe,O, by V(pic),. 0, Measurements made in solution containing 3.5 x mot dm-3 V(pic)-+O.l mol dm-3 HCO; +20 ppm Triton X-100, pH 4.4, T = 80 OC and using nickel ferrite sample Al. Rate constant estimated from amount of Fe dissolved. A, As (O), but rate constant estimated from amount of Ni dissolved. 0, Measurements made in solution containing 3.0 x rnol dmp3 V(pic), +0.1 mol dmp3 CH3CO; +20 ppm Triton X-100, pH 4.4, T = 60 OC and using nickel ferrite sample Al.Lines calculated according to eqn (1 l), using constants shown in table 7. I I I I I I 3.0 3.5 4.0 4.5 5.0 5 . 5 6.0 PH FIG. 7.-Effect of pH on the kinetics of dissolution of NiFe,O, by V(pic);. 0, Measurements made in solutions containing 3.5 x mol dm-3 free picolinate+O.l mol dm-3 HCO;+20 ppm Triton X-100, T = 80 O C and using nickel ferrite sample Al. 0, Measurements made in solution containing 3.0 x mol dm-3 free picolinate+O.l rnol dm-3 CH,CO; + 20 ppm Triton X-100, T = 60 OC and using nickel ferrite sample Al. Lines calculated according to eqn (1 l), using constants shown in table 7. mol dm-3 Vbic); + 3.0 x mol dm-3 V(pic);+ 1.8. x1158 METAL OXIDE DISSOLUTION KINETICS TABLE EFFECT OF BUFFER ON THE v@ic); + NiFe,O, REACTION^ ~ ~~ ~ buffer (kobs/[V"l)b conc/mol dm-3 type pH /dm3 mol-l min-' 0.1 HCO; 4.3 1.6 0.2 HCO; 4.6 1.8 0.5 HCO; 4.9 1.2 0.1 CH3C0, 4.4 2.2 a Measurements made in solutions containing also ca.3 x 2.9 x lo-, mol dm-3 free picohate+ 20 ppm Triton X-100, T = 80 O C ; on the amount of iron dissolved. mol dm-3 V(pic),+ Rate constants based TABLE 3.-EFFECT OF SURFACTANT ON THE v(pic); 4- NiFe,O, REACTION' sur fac tan t (kobs/[V"I) conc (ppm) typeb /dm3 mol-l mind' 1.6 20 T 1.6 100 T 1.5 100 NaLS 1.2 100 H 2.5 - 0 a Measurements made in solution containing also ca. 3 x mol dm-3 V(pic); + 2.9 x T = Triton X-100, lo-, mol dm-3 free picolinate + 0.10 mol dm-3 HCO;, pH 4.4, T = 80 "C; NaLS = Sodium lauryl sulphate, H = Hyamine 1622. SURFACTANT Varying the concentration of the surfactant, Triton X-loo,* had practically no effect on the rate (table 3).Within experimental error there was no change in rate on omitting the surfactant altogether, and only minor changes on substituting an anionic surfactant, sodium lauryl sulphate, or a cationic one, Hyamine 1622.t BLANK EXPERIMENTS, etc. A number of blank experiments were also done to examine the dissolution behaviour in the absence of VII. As the results in table 4 show, none of the othei reagents used in this work bring about dissolution, except V(pic),, the oxidation product of the V1* reagent. This dissolves nickel ferrite at least fifty times more slowly than V(pic),-, and thus makes a negligible contribution to the processes described here. Changes in stirring rate were also without effect.* Triton X-100 is a proprietary name for a non-ionic surfactant composed of iso-octylphen- Hyamine 1622 is a proprietary name for di-isobutylphenoxyethoxyethyldimethylbenzyl ammonium oxypolyethoxyethanol containing cu. 10 ethoxy units. chloride monohydrate.M. G. SEGAL AND R. M. SELLERS 1159 TABLE 4.-EFFECT OF STIRRING RATE AND ADDITIVES ON THE V(piC), 4- NiFe,O, REACTION, AND BLANK EXPERIMENTS reagents, etc. effect of stirring ratea no stirring 1.9 small stirring bar at maximum speedb 1.6 large stirring bar at maximum speed 1.6 effect of other additives" no other additives 0.10 mol dm-3 NaCl 0.10 mol dm-3 Na,SO, 1.6 1.7 1.8 blank experimentsC 0.04 mol dm-3 picolinate + 0.10 mol dm-3 < 1 % dissolution in 4?j h 1.0 mol dm-3 H,SO, < 1% dissolution in h 6.4 x 1 0-3 rnol dm-3 V(pic), + 0.022 mol dm-3 HCO;, pH 4.6 0.033d free picolinate, 0.10 mol dm-3 HCO;, pH 3.9 a Measurements made in solution containing also ca.3 x mol dm-3 mol dm-3 free picolinate + 0.10 mol dmW3 HCO; + 20 ppm Triton X-100, Solutions also V(pic); + 2.9 x pH 4.3, T = 80 OC; * Standard conditions used for all other dissolution runs; contained 20 ppm Triton X-100; T = 80 OC; kOb,/[V1ll]. 4 3 n - I z g2 -3 E + s! - 1 0 2.8 2.9 3 .O 3.1 3.2 I 0 3 KIT FIG. 8.-Arrhenius plots for the dissolution of NiFe,O, by V(pic);. 0, n = 2. Measurements made in solutions containing 3.5 x mol dm-3 V(pic); + 0.030 mol dmb3 free picolinate +O. 1 mol dm-3 HCO;+20 ppm Triton X-100, pH 4.4 and using nickel ferrite sample Al. 0, n = 3.Measurements made in solution containing 3.0 x mol dm-3 V(pic);+O.O18 mol dm-3 free picolinate+O.l mol dm-3 CH3CO; + 20 ppm Triton X-100, pH 4.4 and using nickel ferrite sample Al. Lines calculated from least-squares analysis of data.1160 METAL OXIDE DISSOLUTION KINETICS TABLE 5.-ACTIVATION ENERGIES FOR THE v(piC), + NiFe,O, REACTION buffer EJkJ m o P formate 67+ 10 acetate 72+11 EFFECT OF TEMPERATURE ON THE DISSOLUTION KINETICS The influence of temperature on the dissolution rate was investigated for both formate and acetate buffers in the temperature range 40-80 OC. Good agreement with the Arrhenius law was obtained, as shown in fig. 8. Activation energies estimated from these plots are shown in table 5. The figures are very similar, and the reaction is'clearly independent of the nature of the buffer.TABLE 6.-EFFECT OF OXIDE CALCINING TEMPERATURE ON THE KINETICS OF THE V(pic), + NiFe,O, REACTION^ sample oxide preparation calcining (kobs/[VI'I) no.b method conditions /dm3 mo1-I min-' A1 A2 A3 c 1 c 2 c3 c 4 c 5 C6 oxalate oxalate oxalate carbonate carbonate carbonate carbonate carbonate carbonate 6 h at 1000 OC 12 h at 1000 OC 24 h at 1400 OC 6 h at 600 OC 6 h at 800OC 6 h at 1000 OC 6 h at 1200 OC 6 h at 1400 OC 72 h at 1400 OC 1.6 1.7 0.90 5.0 0.41 0.34 0.31 0.29 ca. 17 Measurements made in solution containing ca. 3 x mol dm-3 V(pic);+ mol dm-3 free picohate+ 0.10 mol dm-3 HCO, +20 ppm Triton X-100, pH 4.3, 2.9 x T = 80 OC; CJ table 1. EFFECT OF OXIDE CALCINING TEMPERATURE ON THE DISSOLUTION KINETICS Results on the effect of calcining temperature on the dissolution of NiFe,O, by V(pic); are given in table 6.The range of temperatures over which calcining could be varied was limited to ca. 1000-1400 OC. The lower limit is set by the need for a temperature of at least 800 OC in order to form the spinel phase, and the upper limit by the materials of construction of the calcining furnace and the oxide itself. The two oxides calcined at 600 and 800 O C (samples C1 and C2, respectively) contain Fe,O,, and this presumably accounts for their much higher dissolution rates. For the oxides prepared by the oxalate method there was a reduction in rate of about a factor of 2 on increasing the calcining temperature from 1000 to 1400 OC but a much smaller change with the 'carbonate' oxides.With neither group of oxides did the length of the calcining period have any effect.M. G. SEGAL A N D R. M. SELLERS 1161 The decreasing trend in rate constant with increasing calcining temperature probably arises from changes in the surface areas of the oxides through sintering. The differences in the relative decrease are more difficult to understand. They may reflect the slightly different calcining histories of the two groups of oxide [see footnote (a), table 11, although this seems unlikely, or it may be that some characteristic of the oxide (its fault structure?) is more easily annealed out or modified at 1000 O C when prepared by the carbonate method than when prepared by the oxalate method. DISCUSSION THE REDUCTIVE DISSOLUTION PROCESS The variation of the dissolution rate with time and its dependence on the size of the particles establish clearly that the dissolution process involves reaction at the particle surface as the rate-determining step, and the linear dependence on [v(pic);] that the species reacting at the surface is V(pic);.For reasons outlined elsewhere12 we ascribe this to an outer-sphere electron-transfer to FeIII ions in the surface (10) V(pic); + >FelIJ -+ V(pic), + >FeII. The other reagent dependences we attribute to the effect on reaction (10) of adsorption of H+, picolinate, etc. at the surface according to a simple model based on the Langmuir adsorption isotherm. Assuming the reactions occurring to be those shown in the following scheme, kobs will vary with reagent concentration according to eqn (1 1).kll V(pic); + >s -+ V(pic), + >s- K b >s+H+’s-H+ k b V(pic); + >s-H+ -+ V(piC), + >s-H K c >s + L * >s-L kc V(pic); + >s-L + V(pic), + x-L- (L = picolinate, without regard to its state of protonation with >s = surface site.) This equation has been used to calculate the lines shown in fig. 5-7, using the values of the constants given in table 7, and accounts well for the results obtained. The changes in the five constants with temperature are much as expected with the rate constants k,, k , and k, increasing, and the binding constants Kb and K, decreasing (K, markedly) with increasing temperature. It must be stressed, however, that the mechanism as represented in the scheme is a simplification. In particular it neglects adsorption of picolinate at protonated sites, and assumes that the anionic and zwitterionic forms of picolinic acid behave identically.This is certainly not the case in the adsorption of picolinate onto haematite, where we find that the binding constant reaches a maximum at pH ca. 4.8.3 Partly as a result of these assumptions it will be1162 METAL OXIDE DISSOLUTION KINETICS TABLE 7.-RATE AND EQUILIBRIUM CONSTANTS FOR SOME OF THE PROCESSES IN THE V(pic); + NiFe,O, REACTION rate equilibrium constanta reaction acetate media, 60 OC formate media, 80 O C ka 0.4 k 0.1 0 . 7 5 f 0.25 kb 6 f 2 13f4 kC 0.10 & 0.03 0.75 k 0.25 Kb 3700 f 1200 2300 f 700 KC 150250 20&7 a k,, kb and kc in units of dm3 mo1-I min-l, Kb and Kc in units of mol dm-3. seen that the variation of kobs with [pigolinate] at 80 OC is attributed solely to the replacement of protonated sites by >s-pic as the picolinate concentration increases.More complex mechanisms could be written, but in view of the relatively large uncertainties in the measured rate constants we do not feel that they would contribute materially to our understanding of the factors determining the dissolution kinetics. To summarise, we conclude that adsorption of H+ increases the rate of the reaction, whereas picolinate has an inhibiting effect in comparison with the ‘free’ surface (i.e. kb > k , > kc), and that protons are more strongly bound to the surface than picolinate The surface sites, >s, can be identified with FeIII ions. Protons probably add to (12) Unfortunately it is not known which crystallographic planes are present at the surface of nickel ferrite particles, but undoubtedly the surface contains more than one type of hydroxide (e.g.bound to one Fe3+, to 2 Fe3+ or to Ni2+), as found with other FeIII-containing oxides such as goethite (a-Fe00H).19 Adsorption of picolinate occurs by displacement of these surface hydroxides. In the scheme given above this is written as occurring at a single site, though this may not be correct, for we have evidence in the adsorption of picolinic acid on haematite that two surface sites are occupied per molecule of ad~orbate.~ Similar behaviour has been found with oxalic acid, selenious acid, orthophosphoric acid and sulphuric acid on goethite, e t ~ . ~ ~ - ~ ~ The Fe3+ ions in nickel ferrite are not all equivalent.The oxide’s structure is that of an inverse spinel2s in which the 0,- ions exist in a cubic close-packed array, with the Ni2+ and half the Fe3+ ions occupying the larger octahedral holes and the remaining Fe3+ ions the smaller tetrahedral holes. Whether these two types of Fe3+ behave differently or not is unknown, and raises the further question of whether V(pic); attacks surface Fe3+ ions indiscriminately (be they in octahedral or tetrahedral holes), or only those at surface defects such as ledges, kinks, dislocations, etc. The kinetic results are equivocal on this point, but the morphological changes indicate some localised attack. Faults in the oxide structure seem to be preferred, but more generalised attack also occurs at a measurable rate.The dissolutions of Fe,03 and Fe30, by V(pic); are so rapid3 in fact, that we suspect they approach the diffusion-controlled limit and hence involve little discrimination between surface sites. This interpretation is not certain, however, for complications arise from the charge on the particles (in our calculations (Kb > &)* the surface hydroxides to give bound water according to reaction (12) >Fe-OH + H+ f >Fe-OH,.M. G. SEGAL AND R. M. SELLERS 1163 we have assumed it to be zero) and the highly porous nature of the oxides used in these experiments. We hope to investigate these reactions in more detail. The processes subsequent to the reduction of the surface Fe3+ ions, i.e. the actual dissolution reactions, are not very clear but presumably involve some kind of terrace-ledge-kink me~hanism.~' The driving force for the disruption of the oxide surface and the ejection of Fe2+ ions into the bulk solution is the increase in size of iron ions on reduction (typical crystallographic radii for the high-spin ions are 0.75 A for Fe2+ and 0.65 A for Fe3+ 28), and the increased electrostatic repulsion between the electron clouds of the iron ions and the adjacent 02- ions.Nickel ions probably pass into solution after removal of most or all of the neighbouring ferric ions, and this probably accounts for the disparity between iron and nickel in solution in the early stages of the dissolution. Some Fe3+ ions may even pass into solution in this way. They would, however, be immediately reduced by V(pic); in the bulk solution, making such a pathway indistinguishable from that involving reduction prior to dissolution.COMPARISON WITH ACID DISSOLUTION REACTIONS An extensive literature exists on the dissolution of metal oxides by acid, including both theoretical descriptions of the processes i n ~ o l v e d , ~ ~ - ~ ~ and experimental studies of the factors that influence the kinetics.l07 32-38 Acid dissolutions occur by attack of protons at defects in the oxide surface and give rise in the main to pitting and other localised forms of attack. This is to be contrasted with V(pic),, where in addition some more general surface attack seems to occur. V(pic); also differs in its sensitivity towards surfactants. Jones et have shown that adsorption of such compounds at nickel oxide surfaces can effectively block dissolution by acid, whereas the experiments described here suggest no dependence on surfactant.The differences probably arise because of the need in acid dissolution for the proton to diffuse right up to the particle surface before reaction can take place, whereas electron transfer can occur ' through' the surfactant, as found in the reaction of the hydrated electron with organic compounds solubilised in m i ~ e l l e s . ~ ~ ~ 40 TECHNOLOGICAL APPLICATIONS The results presented here describe a new and rapid means of dissolving oxide deposits containing FeI" and in particular NiFe204.41 V(pic); and similar reductants [which have been dubbed LOMI (low oxidation-state metal ion) reagents] clearly show much promise for use in the cleaning of oxide deposits in power plant, especially in pressurised-water reactors where the deposited oxide is principally NiFe20,.42 The cleaning of various reactor artefacts by V(pic); and the use of the reagent for a full-scale reactor cleaning have been described el~ewhere.~~ This work was performed in part under contract RP1329-1 with the Electric Power Research Institute, and is published by permission of the Central Electricity Generating Board. We thank Drs D.Bradbury, T. Swan and C. J. Wood for their advice and encouragement, Messrs B. Daniel and G. Marsh for experimental assistance, and Dr P. Tempest for help with the X-ray measurements. D. G. Wickham, Inorg. Synth., 1967, 9, 152. R. Bock, A Handbook of Decomposition Methoh in Analytical Chemistry (International Textbook Co., London, 1979), p.83, and references 4.317-4.322 therein. D. Bradbury, M. G. Segal, R. M. Sellers, T. Swan and C. J. Wood, unpublished results. (a) E. Peligot, Ann. Chim. Phys., 1844, 12, 533; (b) E. Pkligot, Ann. Chim. Phys., 1845, 14, 240.1164 METAL-OXIDE DISSOLUTION KINETICS (a) H. W. Swank and M. G. Mellon, Znd. Eng. Chem., Anal. Ed., 1938, 10; (b) W. Klump and H. Busch, Mitt. Ver. Grosskesselbesitzer, 1962, 81, 433; (c) A. L. Wilson, Analyst, 1964,89, 402; ( d ) J. A. Tetlow and A. L. Wilson, Analyst, 1964, 89, 442. J. A. McKeague and J. H. Day, Can. J. Soil Sci., 1966, 46, 13. T. Tamagawa, Kokai, 1974, 74-14,629. B. A. Zabin and H. Taube, Znorg. Chem., 1964, 3, 963. lo N. Valverde, Ber. Bunsenges. Phys.Chem., 1976, 80, 333. '* D. Bradbury, in Water Chemistry of Nuclear Reactor Systems (British Nuclear Energy Society, l2 M. G. Segal and R. M. Sellers, J . Chem. SOC., Chem. Commun., 1980, 991. l 3 M. J. Pryor and U. R. Evans, J. Chem. SOC., 1950, 1259. l4 S. Haruyama and K. Masamura, Corros. Sci., 1978, 18, 263. l5 D. G. Wickham, E. R. Whipple and E. G. Larson, J. Znorg. Nucl. Chem., 1960, 14, 217. l6 Powder Diffraction File(Joint Committee on Powder Diffraction Standards, Swarthmore, Pennsylvania, l7 Handbook of Chemistry and Physics, ed. C. R. Weast (Chemical Rubber Co., Cleveland, Ohio, 50th l8 (a) R. C. Mercier and M. R. Pans, C.R. Acad. Sci., 1964,259,2445; (b) R. C. Mercier, M. Bonnet and Is J. D. Russell, R. L. Parfitt, A. R. Fraser and V. C. Farmer, Nature (London), 1974, 248, 220.2o R. J. Atkinson, R. L. Pafitt and R. St. C. Smart, J. Chem. SOC., Faraday Trans. I , 1974, 70, 1472. 21 R. L. Parfitt, R. J. Atkinson and R. St. C. Smart, Soil Sci. SOC. Am. Proc., 1975, 39, 837. 22 J. D. Russell, E. Paterson, A. R. Fraser and V. C. Farmer, J. Chem. Soc., Faraday Trans. 2,1975,71, 23 R. L. Padtt, J. D. Russell and V. C. Farmer, J. Chem. SOC., Faraday Trans. I , 1976, 72, 1082. 24 R. L. Parfitt and R. St. C. Smart, J. Chem. SOC., Faraday Trans. I , 1977, 73, 796. 25 R. L. Parfitt, V. C. Farmer and J. D. Russell, J. Soil Sci., 1977, 28, 29. 26 A. F. Wells, Structural Inorganic Chemistry (Clarendon Press, Oxford, 4th edn, 1975), p. 490. 27 G. M. Rosenblatt, in Treatise on Solid State Chemistry, ed. N. B. Hannay (Plenum Press, New York, 28 A. F. Wells, Structural Inorganic Chemistry (Clarendon Press, Oxford, 4th edn, 1975), p. 259. 28 H.-J. Engell, 2. Phys. Chem. (N.F.), 1956, 7 , 158 (CE-trans. 6891). 30 D. A. Vermilyea, J. Electrochem. Soc., 1966, 113, 1067. 31 N. Valverde and C. Wagner, Ber. Bunsenges. Phys. Chem., 1976, 80, 330. 32 J. W. Diggle, in Oxides and Oxide Films, ed. J. W. Diggle (Marcel Dekker, New York, 1973), 33 M. Simnad and R. Smoluchowski, J. Chem. Phys., 1955, 23, 1961. 34 (a) K. Azuma and H. Kametani, Trans. Metall. SOC. AZME, 1964, 230, 853; (6) H. Kametani and 35 I. H. Warren, M. D. Bath, A. P. Prosser and J. T. Armstrong, Znst. Min. Metall. Trans., Sect. C, 1969, 36 R. M. Cornell, A. M. Posner and J. P. Quirk, J. Znorg. Nucl. Chem., 1976, 38, 563. 37 (a) C. F. Jones, R. L. Segall, R. St. C. Smart and P. S. Turner, J. Chem. SOC., Faraday Trans. I , 1977, 73, 1710; (b) C. F. Jones, R. L. Segall, R. St. C. Smart and P. S. Turner, J. Chem. SOC., Faraday Trans. I , 1978, 74, 1615, 1624; (c) R. L. Segall, R. St. C. Smart and P. S. Turner, J. Chem. SOC., Faraday Trans. I , 1978, 74, 2907. ' M. Miyazaki, M. Amemiya, Y. Sat0 and T. Takamura, Kokai, 1972, 72-25,073. London, 1978), p. 373. 1974). edn, 1969), p. B232. M. R. Pans, Bull. SOC. Chim. Fr., 1965, 2926, 3527. 1623. 1976), vol. 6A, p. 165. vol. 2, p. 281. K. Azuma, Trans. Metall. SOC. AZME, 1968, 242, 1025. 78, 21. 38 K. Sangwal, J. Muter. Sci., 1980, 15, 237 and references therein. 38 Th. Proske, Ch-H. Fischer, M. Gratzel and A. Henglein, Ber. Bunsenges. Phys. Chem., 1977,81,816. 40 M. A. J. Rodgers, D. C. Foyt and Z. A. Zimek, Radiat. Res., 1978,75, 296. 41 C. J. Wood, D. Bradbury, T. Swan, M. G. Segal and R. M. Sellers, U.K. Patent Appl. 8,000,584,1980. 42 Y. L. Sandler, Corrosion, 1979, 35, 205. 43 D. Bradbury, M. G. Segal, R. M. Sellers, T. Swan and C. J. Wood, in Water Chemistry of Nuclear Reactor Systems 2 (British Nuclear Energy Society, London, 1981), p. 403. (PAPER 1 /756)

ISSN:0300-9599    

DOI:10.1039/F19827801149

出版商:RSC    

年代:1982

数据来源: RSC