摘要:

J. Chem. SOC., Faraday Trans. 1, 1982, 70, 2497-2508 Absolute Rate Coefficients in the Alternating Copolymerization of Methyl Acrylate and Styrene, Including a Note on the Application of Non-stationary Conditions to Copolymerization BY CLEMENT H. BAMFORD*~' AND PETER J. MALLEY Donnan Laboratories, University of Liverpool, P.O. Box 147, Liverpool L69 3BX Received 26th October, 198 1 A kinetic study has been made of the alternating copolymerization of methyl acrylate (MA) and styrene (St) in the presence of ethylaluminium sesquichloride with photoinitiation (A = 436 nm) by Mn,(CO),, + CCl,. The reaction was carried out in dilute solution in toluene at 25 OC; variation of [St] over a wide range enabled the kinetic parameters kba k&, and kab k& to be evaluated (kba and kab being cross-propagation rate coefficients and ktbb and ktaa termination rate coefficients).Termination coefficients were determined by the rotating-sector method. The application of this technique to copolymerizations is considered and an equation derived for calculating termination coefficients from sector observations. This procedure is valid so long as the coefficients are independent of the reactant composition. The cross-propagation coefficients kba, kab, evaluated from the above data, are found to be large and are compatible with the high degree of alternation observed. It is considered that the simple mechanism of alternation based on dominant cross-propagation is the best representation of the copolymerization. Reasons for the enhancement of propagation coefficients by the presence of Lewis acids are considered.A major influence on kba is attributed to the weakening of the C=C double bond in MA brought about by complexation of the latter with AI,Et,Cl,. We have recently s t ~ d i e d l - ~ thealternatingcopolymerization ofsomevinyl monomers in the presence of the Lewis acid ethylaluminium sesquichloride and have shown that under these conditions the kinetic behaviour of the systems methyl acrylate + ~tyrene,l-~ methyl acrylate + butadiene31 and methyl acrylate + isoprene3$ is consistent with the simplest mechanism of alternation, namely that based on dominating cross- propagation. For this mechanism to be compatible with the high degrees of alternation observed it is clearly necessary for the cross-propagation rate coefficients to be significantly larger than other relevant propagation coefficients in the system.This paper is concerned with the measurement of the former coefficients for the monomer pair methyl acrylate (MA) + styrene (St). Cross-propagation coefficients in copolymerizations which do not show strong alternation are normally evaluated from reactivity ratios and the appropriate homopropagation coefficients. This procedure is unsatisfactory with highly alternating reactions since the reactivity ratios are very small and cannot easily be determined accurately. In such cases the most useful route to the cross-propagation coefficients lies through determination of absolute terminating coefficients and combination of these with other appropriate kinetic parameters.We have adopted this approach in the studies now reported. Hospital, Liverpool L7 8XP. t Present address : Bio-engineering and Medical Physics Unit, Duncan Building, Royal Liverpool 81 2497 FAR 12498 COPOLYMERIZATION OF METHYL ACRYLATE AND STYRENE Few measurements of absolute termination rate coefficients in copolymerizations have been described. This is hardly surprising, since for systems without strong alternation such determinations have seemed unlikely to provide significant new information. We have used the conventional rotating-sector technique. The application to copolymerization of this and other procedures involving non-stationary conditions (e.g. observation of pre- and after-effects) has not been considered in the literature; nevertheless, it differs in some respects from the simple application to homopoly- merization and therefore requires discussion, which is provided later.As in earlier work'. manganese carbonyl [Mn,(CO),,] +carbon tetrachloride was the photo- initiating system, used with light of wavelength 436 nm. We have previously given reasonsly for believing that the rate of photoinitiation obtained is not affected by the nature or concentrations of the monomers present, or the presence of the Lewis acid. In these circumstances, rates of initiation may be calculated from observations of the rate of polymerization of methyl methacrylate, with identical conditions of photoini tiation. The Lewis acid employed was ethylaluminium sesquichloride (Al,Et,Cl,), which, according to Hirooka,6 complexes very strongly with methyl acrylate monomer, the equilibrium constant exceeding 40 mol-l dm3 at 25 OC.Under our conditions (methyl acrylate in excess) effectively the whole of the aluminium derivative would be complexed. We follow Hirooka in denoting the complex by MA. - .a1 (a1 being an aluminium atom); then for our purposes the concentration of complex is given by [MA - - * all = 2[A1,Et3C1,],, where [A1,Et3C1,], is the concentration of ethylaluminium sesquichloride added. EXPERIMENTAL All experiments were carried out in a laboratory illuminated by inactive (sodium) light. MATERIALS Styrene was freed from inhibitor by several washings with 20% aqueous NaOH followed by numerous shakings with distilled water. It was dried overnight by calcium chloride under nitrogen and then fractionated. The monomer was stored at 0 OC over calcium hydride in a nitrogen atmosphere and distilled on the vacuum line immediately before use.Methyl acrylate was purified as described previ~usly.~ Manganese carbonyl was sublimed in vacuum and stored at 0 OC under nitrogen in the dark. Ethylaluminium sesquichloride was purchased as a 25% w/v solution in toluene. It was transferred to ampoules of capacity ca. 5 cm3 for storage under vacuum. This operation, and the subsequent transference of Al,Et,Cl, solution to the reaction vessels, was carried out under a nitrogen blanket. Toluene was dried over sodium wire for at least 24 h before use. Carbon tetrachloride (AnalaR) was used without further purification. When making up reaction mixtures the components were added in the order manganese carbonyl (in toluene), toluene (solvent), carbon tetrachloride (in toluene), ethylaluminium sesquichloride (in toluene), methyl acrylate, styrene.TECHNIQUES Rates of copolymerization were measured at 25 +_ 0.0 1 OC with the aid of dilatometers having cylindrical bulbs 0.5-1 cm3 in capacity and precision-bore capillaries of internal diameter 1 mm. The rate of polymerization was calculated from the relation d dt - -([MA] + [St]) = KR mol dm-, s-lC. H. BAMFORD A N D P. J. MALLEY 2499 where R is the observed rate of contraction in cm per min per unit volume of dilatometer bulb and the constant K has the value K = 8.03 x mol dm-3 s-l/(cm min-l ~ m - ~ ) determined from density measurements on the monomers and alternating copolymer.The optical system comprised a high-pressure mercury arc (Mazda 250 W, type ME/D), used in conjunction with a 436 nm light filter and appropriate glass lenses for obtaining a parallel beam of light. Measurements of 'initial' rates of copolymerization were made as soon as an effectively constant rate of contraction had been established after starting irradiation. Normally this required periods of ca. 1 min, arising from the natural induction period and inhibition attributable to residual traces of impurities, which are virtually impossible to eliminate. Rates of copolymerization in our experiments were large and with high light intensities and low styrene concentrations, significant styrene consumption could occur before the rate became constant.In these circumstances it was found more satisfactory first to illuminate the system by a low intensity (obtained by using a neutral density filter) until a steady rate was achieved. On removing the filter, the constant rate corresponding to the higher intensity could easily be measured. The correction for the styrene consumed in the cleaning-up process was estimated from the overall contraction which had occurred. The rotating-sector equipment was essentially the same as that described by Bamford and Brumby,* except that light filters appropriate for A = 436 nm were used. Copolymers required for analysis were prepared in vacuum in reaction vessels of ca. 10 cm3 capacity. After polymerization the copolymers were precipitated into methanol, reprecipitated, then dried to constant weight.Alternation was examined by elemental analysis and n.m.r. observations on CDCl, solutions of the copolymers.s RESULTS AND DISCUSSION STATIONARY SYSTEMS In earlier w ~ r k l - ~ kinetic observations were limited to ' high' styrene concentrations (2 0.4 mol dm-3); under these conditions the kinetic relation eqn (1) first reported by Hirookasy9 is followed co = k[MA - - - all [MAfree]' [St]" where w is the rate of copolymerization defined by d dt o = --([MA]+[St]). At sufficiently low styrene concentrations eqn (1) does not hold; the dependence of w on [St] over the whole range examined (at constant [MA], [A12Et,C13)] is presented in fig. 1. When the styrene concentration is very low the rate of copolymerization becomes first-order in [St] and appears to be zero-order in [MA - - all.The high rates of copolymerization obtained in these alternating copolymerizations, even at low monomer concentrations, may be appreciated from this figure. Since use of the rotating-sector requires proportionality between co and 94 (9 = rate of chain starting) we have confirmed this for several systems; an example is given in fig. 2. The sharp reduction in the rate of copolymerization as [St] falls below ca. 0.1 mol dm-3 (fig. 1) suggests a dependence of o involving [StI2; in fact, the data fit closely a relation of the form (2) - I+- w2 Ftl2 C 81-22500 COPOLYMERIZATION OF METHYL ACRYLATE AND STYRENE 5 Y I I I I I I I I I I I I 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 1.1 1.2 [ S t ] /mol dm-3 FIG. 1.-Dependence of rate of copolymerization w on [St].Photoinitiation (1 = 436 nm) by Mn,(CO),, + CCl,. Initial concentrations/mol dm-, : MA 0.4, Mn,(CO),, 2.8 x lo-,, CCl, 0.1, Al,Et,Cl, 0.15. Toluene solution, 25 OC. (-) Calculated from eqn (2); (---) calculated from eqn (10); 0, experimental points. 1) (arb. units) FIG. 2.-Dependence of rate of copolymerization w on incident intensity I,. Photoinitiation (A = 436 nm) by Mn,(CO),,+CCl,. Initial concentrations/mol drn-,: MA 0.4, St 0.1, Mn,(CO),, 2.8 x CCl, 0.1, Al2Et,C1, 0.15. Toluene solution, 25 OC. where o, (the limiting value of o at high [St]) and C are constants under present conditions. This is illustrated by fig. 3; note that this plot is very sensitive to errors in o and [St]. A repeat of the experiments with a lower light intensity gave similar results.Experimental values of C and o, are given in table 1 . The solid curve of fig. 1 has been calculated from eqn (2) with the aid of the mean value of C and the appropriate value of o, given in table 1, and it is an excellent representation of the experimental data. Note that in the experiments of fig. 1 the copolymers formed have degrees of polymerization > 1000, even for the lowest [St] used, and they are all very strongly alternating, so that their composition is effectively invariant, irrespective of changes in the reactant concentrations. Furthermore, these latter involve only replacement of solvent (toluene) by styrene or uice uersa since [MA] and [Al,Et,Cl,] remain constant.C. H. BAMFORD AND P. J. MALLEY -4 0 1 2 3 4 5 6 [St]-2/mol-2 dm6 FIG.3.-Plot of eqn (2), using experimental data in fig. 1. TABLE KINETIC PARAMETERS AT 25 "C [Al2Et3Cl3] = 0.15 mol dm-3 250 1 ktaa kba 102- - 1089 105 m0 103 c kba k d b kab k d a /mol dm-3 s-l /mol dm+ s-l /mo12 dm-6 /mol-i dmg s-4 /mol-i dmt s - ~ ktbb(ka) from eqn (2) and (8) 4.1 46.6 5.15 3.84 16.1 5.72 0.37 13.9 5.74 3.81 15.1 6.37 - means : 5.44 3.83 15.6 6.0 from eqn (10) 4. I 46.6 - 4.1a 27.5b We therefore believe that copolymer-medium interactions are closely similar in all experiments of the series and that conformational changes in the propagating radicals, which would lead to variations in diffusion-controlled termination coefficients,1°-12 are minimal. The experimental data may therefore be usefully considered in terms of kinetic expressions derived on the assumption of constant termination coefficients.The simplest mechanism of alternation, which we have advocated in earlier paper~,l-~ is that based on dominant cross-propagation. This is set out in eqn (3), in which A and B represent the two vinyl monomers and I is the initiator. In view of the effectively complete alternation in our copolymers, homopropagation reactions are omitted ; chain transfer is not included I - K l 9 Ri+A --A* Ri+B --B. - A * + B --Be k,, - B - + A --A. k,, - A - +- A*- polymer ktaa - B - + - B- - polymer k,,,. - A * + - B*- polymer ktab (3)2502 This mechanism leads to a stationary rate of copolymerization given by COPOLYMERIZATION OF METHYL ACRYLATE AND STYRENE in which ( 5 ) The termination coefficient kt is defined so that the total rate of termination is k,[B *I2 ; no generality is lost by employing a summary coefficient in this way and no assumptions are made about the rates of the three individual termination steps.For sufficiently large [B], k, approximates to ktbb so that eqn (4) becomes CO = cc), = 2kba[A] (")$ 0 = 2kab[B] ( L)'. (6) ktbb while for very large [A], k, is given by the last term on the right-hand side of eqn (5) and eqn (4) takes the form (7) ktaa If A and B represent complexed methyl acrylate (MA - al) and styrene, respectively, we see that the above conclusions drawn from this mechanism are compatible with the limiting experimental data. Thus, at high [St], eqn (6) predicts a rate of copolymerization which is first-order in [MA - - * all and zero-order in [St] [compare eqn (l)], while at low [St] the situation is reversed, o being first-order in [St] and zero-order in [MA..all. Combination of eqn (4), ( 5 ) and (6) yields the relation expressing the dependence of o on [A] and [B]. At constant [A] we may compare eqn (8) with the empirical equation eqn (2) (B = St). The two expressions are of the same form, except that the cross-termination term containing ktab in eqn (8) is absent from eqn (2). One interpretation of the latter equation is therefore that there exists an unusual extent of chemical control of the termination reaction; although this is generally considered unacceptable, it may, of course, be a feature peculiar to this type of system and attributable to the complexed Lewis acid.Knowing C [eqn (2) and table 13, we may, with the aid of eqn (S), evaluate the parameter (9) a = - __ :::: (k:)' with the results shown in table 1. Combination of values of o, and 9 gives kb, kdb; the figure for this parameter in table 1 is probably more precise than the slightly smaller estimates (for lower [Al,Et3C13]) in an earlier paper.' Finally, from kb, k;ib and a we obtain kabkda (table 1). We now enquire into the applicability of complete diffusion-control of the termination reactions. In this case ktaa = ktab = ktbb = k; and the rate equation (4) assumes the simple formC. H. BAMFORD A N D P. J. MALLEY 2503 in which, as already explained, we may expect ki to be effectively independent of [B] (= [St]) in our system, for constant [MA] and [A12Et3C13].Eqn (10) contains two adjustable parameters, viz. k,,ki-8 and k,,k:-+. It does not appear to fit the experimental data as closely as does eqn (2); from the values of the above parameters in table 1 we may calculate the broken curve of fig. 1 which probably represents the best fit. Note that although k,, kdb and k,, ki-4 have rather similar values, kab ki-4 is appreciably larger than k,, kc:,. This is immediately evident from fig. 1, in which curve (2) has a greater slope than curve (1) near the origin; these slopes are determined by k,,k;l and krtbk;2,, respectively. Inspection of fig. 1 suggests that the value of k,, k;-* in table 1 is the maximum acceptable. NO N-S T A T ION AR Y SYSTEMS We now examine the information obtainable from non-stationary conditions, for example by the rotating-sector technique. If [Z-] is the total radical concentration [k, being given by eqn (5)].In the stationary state the relative concentrations of -A* and N-B* radicals are given by For long chains the ratio [A*]/[B*] will be effectively maintained in the non-stationary reaction, hence, for all conditions [Z*] = [A*]+[B*] = A[B*] (13) where From eqn (1 1) and (1 3) we see that in which k, = kt/A2. (16) Eqn (15) is identical with that holding for the polymerization of a single vinyl monomer with only one type of propagating radical. It follows that calculations of mean propagating radical concentrations [z ] in non-stationary phases of the copolymerization (required in the rotating-sector treatment) may be taken over without change from those for a single monomer, with k , as the sole termination coefficient.Since the rate of copolymerization in our system is proportional to [Z-] for constant [A], [B] we may therefore use the familiar ‘sector curve’ to evaluate k,, which is given by eqn (16). The termination coefficient measured by the sector technique therefore depends not only on the three termination coefficients in kt [eqn (5)] but also on the propagation coefficients k,,, k,, through eqn (14). The physical significance of this result may be appreciated by considering a simplified case in which ktab = ktaa = 0, so that there is only one termination coefficient k, = ktbb. At first sight it may appear surprising that ktbb/A2 rather than ktbb is obtained from sector experiments. At the beginning of a dark period, -Be radicals decay by mutual termination, but at the same time some are regenerated by2504 COPOLYMERIZATION OF METHYL ACRYLATE AND STYRENE the propagation *A* + B, to maintain the ratio [A*]/[B*] effectively constant at the value given by eqn (12).The overall rate of decay of [B-] is therefore smaller than that arising solely from mutual interaction, and corresponds to the effective termination coefficient kt/R, which is .c k,. Similarly, the build-up in [Be] at the start of a light period is slower by a factor R than would be the case in the absence of propagation steps. Due to the intervention of the latter, the -A* radicals act as a reservoir of -Be radicals which buffers changes in their concentration. Note that if all three termination steps have the same rate coefficient ki, as in conventional diffusion-control, k, [eqn (5)] reduces to ki( 1 + kba[A]/kab[B])2 or k;12, so that kz = ki.From eqn (16), with the aid of eqn (5) or (14) and the substitution ( ktaa)i kba[Al - ~ 2 - 1 ktbb kab[Bl CB1 v = - -- we obtain the relation 2(v2+2v4+l) ktbb k, = [ 1 +v(2)7 in which 4 is the familiar factor ktab/(ktaa ktbb)4. Rearrangement of eqn (18), after taking the square root, yields l--+-r-= (19) If k, is measured for a series of values of [A]/[B] (i.e. over a rangelof v / ) a plot of (1 +2y14+ ty2)i/ki against v/ should give a straight line of slope kza and intercept k;jb. [Note that I,Y may be calculated with the aid of eqn (8), (9) and (17)]. Such a plot may therefore be used to evaluate ktaa and ktbb and hence [from a knowledge of 4 or from eqn (S)] ktab.Two points about this procedure must be stressed. First, it is valid only if the termination coefficients are effectively constant under the prevailing conditions, or, in general, are independent of [A]/[B]. We have already given reasons for believing that this is true in the present system, conformational changes in the propagating radicals, which would influence termination coefficients, being small. Secondly, although eqn (1 9) has been derived for an alternating copolymeri- zation, it is generally valid for copolymerisation, subject to the conditions mentioned. h b b h a a If y in eqn (19) is replaced by v/’ = l / v we obtain the alternative relation (20) With the aid of eqn (20), ktaa can be obtained from an intercept, which, depending on the quality of the experimental data, may be better defined than the slope in eqn (19)- Fig.4 presents the results of a typical set of rotating-sector experiments. Rates of initiation were calculated from the rates of copolymerization under continuous irradiation for each point; the method of plotting shown in fig. 4, involving Yi explicitly, minimises errors arising from changes in the light intensity. Values of k, so determined are collected in table 2, together with ry and v/’ calculated from eqn (16) and (17). A plot of eqn (20) (4 = 0) which is a more appropriate equation than eqn (19), is shown in fig. 5 and leads to values of ktaa, ktbb presented in table 3. This table also includes k,, and k,, estimated from the termination coefficients and the appropriate parameters in table I.C.H. BAMFORD AND P. J. MALLEY 2505 FIG. 4.-Typical set of rotating-sector experiments. Initial concentrations: St 0.4 mol dm-3, others as for fig. 1. Abscissa scales: lower, experimental data (O), upper, calculated curve. Asymptotes: (-----). t , / s is duration of light period; 6, w, are the mean rate in interrupted light and the steady rate, respectively. TABLE 2.-vALUES OF k, FROM ROTATING-SECTOR EXPERIMENTS [MA] = 0.4 mol dm-3, [A12Et3C1,] = 0.15 mol dm-3 [Stl 1 O-sk,/ /mol dm-3 v/ y' mol-l dm3 .s-l 0.06 18 1.185 0.844 1.58 0.076 0.967 1.034 5.6 0.40 0.184 5.43 6.0 1.2 0.06 1 16.39 4.0 On the basis of complete diffusion-control of the three termination reactions, the best value of k; is given by the mean of the experimental k , figures (since kl = k,) i.e.4.3 x lo6 mol-l dm3 s-l. This hypothesis leads to the values of k,, and k,, presented in table 3. ABSOLUTE RATE COEFFICIENTS The estimated cross-propagation coefficients evidently depend to some extent on whether the calculations are based on the empirical expression eqn (2) or the hypothesis of complete diffusion-control of termination. It is our view that neither procedure can be rejected at present. Fortunately, the uncertainties so introduced do not affect the general conclusions. Both cross-propagation coefficients are relatively large (table 3) and can readly account for the experimentally observed high degrees of alternation. Thus, taking the homopropagation coefficients at 25 O C as 592 l3 and 44 l4 mol-l dm3 s-' for MA and St, respectively, and values of k,, and kab derived from eqn (2) and (8) (table 3) we may calculate the reactivity ratios as2506 COPOLYMERIZATION OF METHYL ACRYLATE AND STYRENE 9 8 0 I 1 I I I I I I I 0 2 4 6 8 10 12 11, 16 *' FIG.5.-Plot of eqn (20) from data in table 2. TABLE ABSOLUTE RATE COEFFICIENTS IN ALTERNATING COPOLYMERIZATION OF METHYL ACRYLATE AND STYRENE AT 25 OC [AI2Et3Cl3] = 0.15 mol dm-3 coefficient vaIue/mol-l dm3 s-l source ktaa 6.3 x lo6 ktbb 6.0 x lo6 kba 9 382 8 502 kab 39 156 57 025 ktab 0 kl 4.3 x 106 fig. 5 eqn (8) fig. 5 mean k, eqn (2) and (8) eqn (10) eqn (2) and (8) eqn (10) Currently, there is no quantitative information about the influence of A1,Et3Cl3 on the homopropagations, although we have indications that k,, may be lower than quoted above under present conditions ([Al,Et3C1] = 0.15 mol dm-3). Thus the figure for rA in eqn (21) is probably a maximum.From eqn (21) we see that the fraction of alternating triads in an equimolar MA/St copolymer is indicating a remarkably high degree of alternation. In an earlier paper1 we have reported that the parameter kfbk;Jb has a value 0.18 mol-t dm: s*, kfb being the absolute coefficient for the chain-transfer reaction between m B o radicals and carbon tetrabromide. With ktbb = 6 x lo6 mol-l dmf s-l we therefore find kfb = 441 mo1-l dm3 S-l. (22)C. H. BAMFORD AND P. J. MALLEY The uncomplexed radical (I) (illustrated below) has a transfer +6 -MA-St- -MA- St* -MA-St . * . ./ a1 a1 -6 2507 constant for reaction with CBr, at 6OoC of 302:15 assuming this holds at 25 OC we estimate that for (I), kfb = 1.33 x lo4 mol-1 dm3 s-l, a value approximately 30 times as large as that given in eqn (22). We therefore conclude that under the conditions of the transfer experiments ([Al,Et3C13] = 0.075 mol dm-3) not more than 1/30 of the -B*-type radicals have the simple structure (I), the great majority being complexed with the Lewis acid, e.g.as in (11) or (111). A reduction in transfer activity towards CBr, accompanying radical complexation with zinc chloride has been reported for systems containing methyl methacrylate by Lachinov et These effects may indicate a reduction in unpaired spin density on the terminal carbon atom, arising from powerful electron-attraction by the Lewis acid.An inductive effect through two saturated carbon atoms could be operative in (11) but would be expected to be small. On the other hand, in (111) the unpaired spin interacts directly (or perhaps through the aromatic ring) with the metal atom, and the effect could be much more significant. Hirookal' has shown that complexation between the alternating MA/St copolymer and Al,Et3C13 is relatively weak; although no quantitative data are yet available, his work suggests that the strong complexation of the -B*-type radical noted above must be attributable in part to involvement of the unpaired spin. In our view, these considerations provide good reasons for believing that (111) is a plausible representation of the structure of the -B* radicals in the presence of A1,Et3C13.An additional point is that the complexation of (I) produces a radical with a bulky penultimate group and this could be responsible for some reduction in transfer activity towards carbon tetrabromide.'. l8 Further work with a different type of transfer agent (e.g. a thioP) is required to elucidate this matter. The cross-propagation coefficient k,, is increased by a factor > 140 by the presence of Lewis acid (table 3: the rate coefficient for the -St- +MA propagation with uncomplexed species is 59 mol-l dm3 s-l).139 l9 Since complexation increases the cationic character of both reactants, we conclude that the weakening of the carbon- carbon double bond by complexation, implied in structure (IV) f 6 - 6 CH22CHZC'0. . . a1 I (IV) is the predominating influence leading to the high rate coefficient in the presence of Al,Et,Cl,.In other words, complexation converts MA into a new highly reactive monomer. The observed enhancement in k,, would correspond to a reduction in activation energy of approximately 3 kcal mol-1 (ca. 12.5 kJ mol-l) if wholly attributable to this. The value of k,, is increased from 3289 mol-1 dm3 s-ll49 l9 b Ya factor exceeding 10 (table 3) in the presence of the Lewis acid. In this case complexation of the methyl acrylate radical would introduce polarity changes favouring propagation. OMe2508 COPOLYMERIZATION OF METHYL ACRYLATE A N D STYRENE Note that in the three systems studied (methyl acrylate with styrene, b ~ t a d i e n e ~ ? ~ or isoprene3~ 5, the slower of the two cross-propagation reactions (i.e.hydrocarbon radical + MA) is subject to greater enhancement by the Lewis acid, which thereby brings about conditions favouring alternation. One of us (P. J. M.) thanks the S.E.R.C. for financial support. We are indebted to Dr M. Hirooka for invaluable help and advice in the early stages of this work. C. H. Bamford, S. N. Basahel and P. J. Malley, Pure Appl. Chem., 1980, 52, 1837. C. H. Bamford and P. J. Malley, J. Polymer Sci., Polymer Lett Ed., 1981, 19, 239. C. H. Bamford, X-Z. Han and P. J. Malley, Plaste Kautsch, in press. C. H. Bamford and X-Z. Han, J. Chem. SOC., Faraday Trans. 1, 1982, 78, 855. C. H. Bamford and X-Z. Han, J. Chem. SOC., Faraday Trans. 1, 1982, 78, 869. M. Hirooka, Doctoral Thesis, Studies on Alternating Copolymerization with Alkylaluminium Halide (Kyoto University, 1971). ' C. H. Bamford and S. N. Basahel, J. Chem. SOC., Faraday Trans. 1, 1978, 74, 1020. * C. H. Bamford and S. Brumby, Makromol. Chem., 1967, 105, 122. M. Hirooka, 23rd IUPAC Congress, Boston, 1971; Macromol. Preprints, 1, 31 1. lo J. N. Atherton and A. M. North, Trans. Faraday SOC., 1962, 58, 2049. l1 A. M. North and D. Postlethwaite, Polymer, 1964, 5, 237. l2 A. M. North, in Reactivity, Mechanism and Structure in Polymer Chemistry, ed. A. D. Jenkins and l3 M. S. Matheson, E. E. Auer, E. B. Bevilacqua and E. J. Hart, J. Am. Chem. SOC., 1951, 73, 5395. l4 M. S. Matheson, E. E. Auer, E. B. Bevilacqua and E. J. Hart, J. Am. Chem. SOC., 1951, 73, 1700. l5 C. H. Bamford and S. N. Basahel, J. Chem. SOC., Faraday Trans. 1, 1980, 76, 1 12. l6 M. B. Lachinov, T. R. Aslamazova, V. P. Zubov and V. A. Kabanov, Vysokomol. Soedin., Ser. A , A. Ledwith (John Wiley, London, 1974), chap. 5. 1975, 17, 1146. M. Hirooka, personal communication. F. M. Lewis, C. Walling, W. Cummings, E. R. Briggs and F. R. Mayo, J. Am. Chem. SOC., 1948,70, 1519. l8 C. H. Bamford and S. N. Basahel, J. Chem. SOC., Faraday Trans. 1, 1980, 76, 107, 1 12. (PAPER 1 / 1666)

ISSN:0300-9599    

DOI:10.1039/F19827802497

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1982, 78, 2509-2520 Ethane and Propane Hydrogenolysis on Ru Catalysts BY SIGNORINO GALVAGNO* Istituto CNR, Trasformazione e Accumulo Energia, Via Salita S. Lucia sopra Contesse 39, 9801 3 Pistunina, Messina, Italy AND JOHANNES SCHWANK University of Michigan, Department of Chemical Engineering, Ann Arbor, Michigan 48109, U.S.A. AND GIUSEPPE GUBITOSA Istituto ‘G. Donegani’, Novara Research Centre, Via G. Fauser 4, 28100 Novara, Italy AND GIORGIO R. TAUSZIK Montedipe S.p.A., Bollate Research Centre, Via S. Pietro 50, 20021 Bollate, Milano, Italy Received 3rd November, 198 1 The hydrogenolysis of ethane and propane has been investigated on ruthenium catalysts and a detailed kinetic study is reported. Activation energies ranging from 70 to 170 kJ mol-I, positive orders of reaction with respect to the hydrocarbon and negative orders with respect to hydrogen were generally found.No direct correlation between the catalytic activity and the metal particle size seems to exist. However, the apparent activation energies and pre-exponential factors increased with increasing ruthenium dispersion. This agrees with the concept that hydrogenolysis reactions are structure-sensitive reactions. Hydrogenolysis reactions of hydrocarbons on metal catalysts have been studied extensively. It is generally accepted that the hydrogenolysis of ethane and propane is a structure-sensitive reaction. Different reaction mechanisms and intermediates have been proposed and extensive reviews exist in the literature.l? The demanding nature of these reactions makes them useful for probing catalytically active surface sites, especially in the case of bimetallic systems consisting of both an active and an inactive metal component.A typical example is a combination of a Group VIII metal, which is active for the hydrogenolysis, and an inactive Group IB metal. This approach has been used previously to investigate bimetallic Ru-Au catalyst^.^? Striking differences in the behaviour of the catalysts were observed in terms of surface composition, bimetallic cluster formation and catalytic activity, by using either SiO, or MgO as support. It was also found that monometallic Ru/SiO, behaved differently compared with R u / M ~ O . ~ - ~ Therefore it seemed important to perform an extensive investigation on several ruthenium catalysts with the objective of exploring whether parameters such as metal-support interactions, preparation method and metal dispersion can influence the hydrogenolysis activity of ruthenium.25092510 HYDROGENOLYSIS ON RU EXPERIMENTAL Supported ruthenium was prepared according to the following methods. (a) The support was impregnated (by the incipient wetness method for all supports except MgO, which was soaked in an excess volume of liquid) with a solution of RuCl, - H,O (Rudi- Pont reagent grade) followed by drying at 110 OC for 4 h. The catalysts were then reduced in flowing hydrogen at 400 OC for 2 h. (b) A solution of Ru [(CH,-CO), CHI, in toluene was heated with the support under reflux for 2 h. The toluene was removed by filtering and the catalyst precursor was washed with toluene until the washings were clear.The red solid obtained was dried under vacuum at 110 OC and reduced in flowing H, at 400 OC for 2 h. (c) Ru,(CO),, in toluene was treated with the support as in method (6). The resulting greyish-brown solid was dried at 110 OC under vacuum and reduced in flowing H, at 400 OC for 2 h. The supports were commercial SiO, (surface area 680 m2 g), A1,0, (surface area $1 60 m2 g-l) and MgO (surface area 15 m2 g-l). The MgO-CaO support (surface area 94 m2 g-l) was prepared by coprecipitation of MgCl, and CaCl, by (NH,), CO,, followed by decomposition at 600 O C . The unsupported ruthenium was a commercial Ru sponge (Baker). A summary of the Ru catalysts studied is reported in table 1.More details concerning the physico-chemical characterization of the samples 3.0 wt Ru/SiO,, 2.9 wt% Ru/A1,0, and 3.1 wt % Ru/SiO,-Al,O, have been previously reported.s The characterization of the samples 4.44 wt % Ru/MgO and 3.86 wt % Ru/SiO, are reported in ref. (3) and (7). TABLE 1 .--PREPARATION AND CHARACTERIZATION OF SUPPORTED Ru CATALYSTS Ru particle size Ru exposed preparation (wt %) support (chemisorption)/ A (%) method 3.86 3.00 2.9 1.16 0.69 3.1 4.44 5.38 4.13 1 00 SiO, SiO, A1203 A1203 SiO,-Al,O, MgO MgO CaO-MgO - 34 22 34 14 13 28 129 16 21 1000 26 43 26 66 72 33 60 44 7.1 0.09 0, chemisorption experiments to determine the percentage of Ru exposed on the surface were performed at room temperature in a flow system using the pulse technique with a thermal conductivity detector.The surface area of the unsupported Ru sample was determined by H, chemisorption at room temperature in a static system. The stoichiometries of chemisorption used were: Ru/O = 4 and Ru/H = 1. The rate of ethane and propane hydrogenolysis was followed in a flow system employing a tubular reactor at atmospheric pressure using helium as diluent. Prepurified hydrogen was passed through Pd asbestos (400 "C) and a molecular sieve trap at liquid-nitrogen temperature. Ultra-high purity He was passed through an oxytrap (Alltech) at room temperature followed by a molecular sieve trap at liquid-nitrogen temperature. Ethane and propane (CP grade) were used without further purification. The reactor, of Pyrex glass, was filled with 50-500mg of catalyst diluted with 0.3 g of ground Pyrex glass. The reactant mixture was fed to the reactor after passing a preheating section. The reactants and products were analysed by a gass.GALVAGNO, J. SCHWANK, G. GUBITOSA AND G. R. TAUSZIK 251 1 chromatograph (HP model 5750 with flame detector) connected directly to the flow system and employing a column filled with silica gel (100-120 mesh). Since preliminary runs showed a decrease of activity with time the following procedure was used to measure the initial rates. The reactant gases were passed over the catalyst for 2 min prior to sampling the products for analysis. The hydrocarbon and helium were then cut out and the hydrogen flow continued for 15 min prior to another reaction period. After 4 or 5 runs the catalyst was treated at 35OOC in flowing H, for 15min and cooled at the reaction temperature in H, before another series of experiments was undertaken.Reaction rates were determined at a partial pressure of hydrocarbon between 1 and 10 kPa and of hydrogen between 10 and 50 kPa. The temperature was varied between 120 and 260 O C . Within the range of flow rates investigated the reaction rates were found to be independent of the gas flow rate, indicating the absence of external mass transfer limitations. Absence of diffusional limitations within the catalyst pellets was verified by calculations based on the criteria developed by Weisz.8 The conversion was kept low (< 5%) throughout the experiments in order to operate under differential conditions. RESULTS PROPANE HYDROGENOLYSIS Rates, V , were calculated from the expression: V/molecule s-l (Ru surface atom)-l = (F/A,)a where F is the feed rate of propane or ethane, A , is the number of Ru surface atoms (by chemisorption) and a is the fraction of consumed hydrocarbon.The catalyst pretreatment, before kinetic measurements, included an in situ reduction in flowing hydrogen at 400 OC for 2 h. Under our experimental conditions, the propane hydrogenolysis led to a products ratio CH,/C,H, > 1 . This indicates that two reactions took place: C3H,+2 H, -+ 3 CH, (1) (2) C3H, + H, -+ CH, + C,H,. The selectivity to C,H, was calculated by the expression: where Vpl and Vp2 are the rates of propane hydrogenolysis according to reactions (1) and (2), respectively. The influence of temperature on the reaction rates was studied at a partial pressure of H, of 20.0 kPa and of propane of 3.0 kPa.The plot of log V against 1/T of all but two samples (namely, 4.44 wt % Ru/MgO and unsupported Ru) gave a straight line for Vpl, V,, and V, (V, is the rate of the total disappearance of propane). A typical plot is reported in fig. 1. From the slope of this and other similar plots the apparent activation energies have been calculated and are reported in table 2. The activation energy of reaction (1) is always higher than that of reaction (2), resulting in a general decrease in the selectivity S with temperature. Samples 4.44 wt % Ru/MgO and unsupported Ru behaved in a quite different manner. Fig. 2 and 3 show the effect of the temperature on the reaction rates for these two samples.For the total disappearance of propane an activation energy of 75 and 88 kJ mol-1 was found on 4.44 wt % Ru/MgO and unsupported Ru, respectively. These values are significantly lower than those for the other samples. Furthermore, at temperatures > 145-150 O C2512 HYDROGENOLYSIS ON RU I I 2.3 2 . L 2.5 103 K I T FIG. 1 .-Arrhenius plot for propane hydrogenolysis on the 3.86 wt % Ru/SiO, sample. pH, = 20.0 kPa, PC,H, = 3.0 kPa; 0, vp1; A, vp2; 0, vp. I I I 2.2 2.3 2.4 103 K I T FIG. 2.-Arrhenius plot for propane hydrogenolysis on the 4.44 wt % Ru/MgO sample. pH, = 20.0 kPa,pcJH8 = 3.0kPa; 0, v,,; 0, vP2; A, vP. the amount of ethane formed on 4.44 wt % Ru/MgO and on unsupported Ru levelled Off. The rates of reaction and the selectivity for all catalysts are compared at 160 O C in table 2.The Si0,-supported samples had, in terms of the overall reaction rate, Vp, theS. GALVAGNO, J. SCHWANK, G. GUBITOSA A N D G. R. TAUSZIK 2513 TABLE 2.-ACTIVATION ENERGY, E,, CATALYTIC ACTIVITY, v, AND SELECTIVITY, s, FOR THE HYDROGENOLYSIS OF PROPANE ON SUPPORTED R U CATALYSTS Vat 160 "C Ru /molecule s-l S at Ru exposed E temperature (Ru surface 160 "C (wt %) support (%) reactiona /kJ mol-l range/"C atom)-' (%) - 3.86 3.00 1.16 0.69 2.90 3.10 4.13 4.44 100 SiO, SiO, 4 0 3 A1203 SiO ,-A1 ,O CaO-MgO MgO - 26 43 66 72 26 33 44 7.1 0.09 1 196 2 109 T 113 1 167 2 125 T 125 1 184 2 130 T 142 1 20 1 2 134 T 134 1 159 2 105 T 104 1 188 2 109 T 109 1 180 2 130 T 142 1 2 T 75 1 2 T 88 - - - - 119-157 13 1-148 173-229 150-188 126-180 151-212 153-195 145-200 120-200 2.8 x lo-, 18.9 x lo-, 21.7 x lo-, 2.1 x 10-2 25.9 x lov2 28.0 x lo-, 1.3 x 10-4 3.3 x 10-3 3.4 x 10-3 1.0 x 10-4 4.4 x 10-3 4.5 x 10-3 5.3 x 10-4 3.1 x 10-3 3.6 x 10-3 1.4 x 10-5 1.5 x 10-3 1.5 x 10-3 1.1 x 10-3 8 .2 ~ 10-3 9.3 x 10-3 3.3 x 10-3 2.0x 10-3 5.3 x 10-3 5 . 2 ~ 10-3 4.5 x 10-3 9.7 x 10-3 87 93 96 99 85 99 88 38 47 a T = reaction (1) +reaction (2). highest activities followed by unsupported ruthenium, ruthenium supported on CaO-MgO, MgO and Al,O, and finally by SiO,/Al,O,, which was the least active. Under the experimental conditions used, C,H, was the main product in the hydrogenolysis of propane, with the exception of 4.44 wt % Ru/MgO and unsupported Ru, on which the selectivity dropped to a value < 50% at 160 O C .After determining the influence of temperature, a value of Tintermediate in the range examined was used to study the effect of propane and hydrogen partial pressure on the reaction rates. The dependence of the rates of reaction on the partial pressure of the reactants can be expressed in the form of a simple power rate low: V = kpgp?. (4) Kinetic orders n and m, calculated from the slopes of the curves of log V against logp, are reported in table 3. On 4.44 wt % Ru/MgO and unsupported Ru the reaction orders were measured at different temperatures and the results are collected in table 4.2514 HYDROGENOLYSIS O N Ru TABLE PR PRESSURE DEPENDENCE EXPONENTS FOR THE HYDROGENOLYSIS OF PROPANE Ru reaction reaction order (wt %) support reactiona order in H,b in propaneC T/OC 3.86 SiO, 1 - 2.5 0.40 143 2 - 1.3 0.67 T - 1.3 0.65 3.00 SiO, 1 - 2.9 0.42 131 2 T 2 T 2 T 2.90 1 2 T 2 T 2 T 1.16 A1203 1 0.69 A1203 1 3.10 Si0,-Al,03 1 4.13 CaO-MgO 1 - 1.5 - 1.6 - 1.7 - 0.7 - 0.9 - 2.9 - 1.8 - 1.8 - 2.4 - 1.2 - 1.3 - 1.9 - 1.5 - 1.5 - 2.7 - 1.4 - 1.6 0.49 0.49 0.50 0.68 0.65 0.48 0.62 0.61 0.57 0.88 0.88 0.54 0.69 0.69 0.81 0.91 0.89 91 79 45 84 69 4.44 MgO 1 - 1.74 .0.74 154 2 + 0.05 1.12 T - 0.82 0.92 1 - 2.2 0.76 154 2 0 0.79 T -0.79 0.80 100 - a T = reaction (1) +reaction (2). Propane partial pressure = 3.0 kPa. Hydrogen partial pressure = 20.0 kPa. ETHANE HYDROGENOLYSIS The reaction of ethane and hydrogen produced methane, according to C,H, + H, -+ 2 CH,. The temperature dependence of the reaction rates was studied under conditions identical to those used for the hydrogenolysis ofpropane (H, partial pressure = 20 kPa, ethane partial pressure = 3 kPa).Fig. 4 shows the temperature dependence of the reaction rates on the different ruthenium catalysts. Table 5 summarizes the apparent activation energies and the activities measured at 160 OC for ethane hydrogenolysis on the Ru catalysts. The Si0,-supported samples had the highest activities, followed by the ruthenium sponge, ruthenium supported on MgO, CaO-MgO, Al,O, and finally Si0,-Al,O,-supported ruthenium, which was clearly the least active. Once again both the unsupported Ru and 4.44 wt % Ru/MgO had a lower apparent activation energy than the other samples. The orders of reaction with respect to both hydrogen and ethane are reported ins. GALVAGNO, J.SCHWANK, G. GUBITOSA AND G. R. TAUSZIK 2515 TABLE PRESSURE DEPENDENCE EXPONENTS FOR THE HYDROGENOLYSIS OF PROPANE ON 4.44 Wt % RU/M@ AND UNSUPPORTED RUTHENIUM AT DIFFERENT TEMPERATURES Ru reaction order reaction order (wt %) support reactiona in H,* in propaneC T/OC 4.44 MgO 1 2 T 1 2 T 1 2 T 1 2 T 1 2 T 1 2 T 100 - 1.74 + 0.05 -0.82 - 1.25 + 0.40 - 0.80 - 0.68 + 0.45 - 0.57 - 2.90 - 0.75 - 1.15 - 2.20 0 - 0.79 - 1.30 +0.41 -0.77 0.74 154 1.12 0.92 0.96 172 1.10 0.99 0.94 194 1.10 0.97 0.81 134 0.98 0.96 0.76 154 0.79 0.80 0.95 177 0.90 0.94 a T = reaction (l)+reaction (2). Propane partial pressure = 3.0 kPa. Hydrogen partial pressure = 20.0 kPa. lo3 KIT FIG. 3.-Arrhenius plot for propane hydrogenolysis on unsupported Ru.pH, = 20.0 kPa, pC,H, = 3.0 kPa; 0, vp~; 0, vp2; A, vp-2516 HYDROGENOLYSIS ON RU -1 .o I I I 2.0 2.1 2 . 2 2.3 1 0 3 KIT FIG. 4.-Arrhenius plot for ethane hydrogenolysis on Ru catalysts. pH, = 20.0 kPa, pC,H, = 3.0 kPa. 3.86 wt % Ru/SiO,, 0 ; 3.0 wt % Ru/SiO,, 0 ; unsupported Ru, 0; 1.16 wt % Ru/A1,0,, 0 ; 4.44 wt % Ru/MgO, A; 0.69 wt % Ru/AI,O,, Is; 4.13 wt % Ru/CaO-MgO, v; 2.9 wt % Ru/AI,O,, @; 5.38 wt % Ru/MgO, I; 3.1 wt % Ru/SiO,-Al,O,, 0. TABLE 5.-KINETE PARAMETERS FOR THE HYDROGENOLYSIS OF ETHANE ON RU CATALYSTS Vat 16OOC /molecule Ru s-l (Ru reaction reaction Ru exposed E, temperature surface order order (wt %) support (%) /kJ mol-l range/OC atom)-l in H2" in ethaneb T/OCC 3.86 3 .oo 1.16 0.69 2.90 3.10 4.13 4.44 5.38 100 SiO, SiO, A1203 A1203 A1203 Si0,-A1203 CaO-MgO MgO MgO - 26 43 66 72 26 33 44 17.1 60 0.09 125 160-191 142 161-196 155 183-214 142 179-213 130 184-217 138 179-247 167 159-200 88 161-225 171 169-200 88 192-243 2.7 x 10-3 1.2 x 10-3 1.6 x 10-4 9.8 x 10-5 2.3 x 10-4 6.1 x 10-5 3.9 x 10-4 7 .7 ~ 10-5 8.3 x 10-4 9.0 x -2.21 - 1.92 -2.37 - 1.45 - 1.12 - 1.45 - 1.55 - 0.73 - 1.93 - 1.10 0.66 160 1.00 186 0.85 206 0.82 190 1.04 203 0.88 242 0.88 180 0.98 189 0.84 185 1.03 188 a Ethane partial pressure = 3.0 kPa. Hydrogen partial pressure = 20.0 kPa. Tem- perature used in determining the order of reaction. table 5. The order of reaction with respect to ethane was always positive and close to unity. The hydrogen orders were always negative, ranging from - 2.4 to - 0.7. This latter value applied for the 4.44 wt % Ru/MgO sample.The catalytic activity measured on our Ru/SiO, catalysts is in excellent agreement with that reported by Sinfelt under similar reaction condition^.^s. GALVAGNO, J. SCHWANK, G. GUBITOSA AND G. R. TAUSZIK 2517 150 - I I 0 E m y: . 3 Q 100 Ru exposed (%) FIG. 5.-Activation energy for overall propane hydrogenolysis as a function of ruthenium dispersion. I I I I 20 40 60 E Ru exposed ( V ) FIG. 6.-Activation energy for ethane hydrogenolysis as a function of ruthenium dispersion. DISCUSSION From the results obtained for ruthenium samples it appears that the same overall behaviour applies for the hydrogenolysis of ethane and propane. Activation energies ranging from 70 to 170 kJ mol-l, positive orders with respect to the hydrocarbon and negative orders with respect to hydrogen were generally found.Furthermore, tables 2 and 5 show that for both reactions the catalytic activity per Ru surface atom measured at 160 OC is higher on Ru/SiO, than on all the other catalysts, with Ru/SiO,-Al,O, being the least active sample. However, due to the large variation in the apparent activation energies and orders of reaction, any ranking of activity is only meaningful for a certain reaction condition. Changing hydrogen partial pressures2518 HYDROGENOLYSIS ON RU 40 30 ‘r: c - 20 I I 100 150 E,/kJ rnol-I FIG. 7.-Plot of pre-exponential factors against activation energies for hydrogenolysis of ethane. and/or temperature can lead to a different picture. No direct correlation between the catalytic activity and dispersion of Ru seems to exist.This raises the possibility that an interaction between the metal and the support modifies the catalytic properties of Ru. However, there seems to be a trend for the apparent activation energies to increase with increasing dispersion (fig. 5 and 6). Similar changes in activation energies for the hydrogenolysis reaction have been observed previously. Calculating the fraction of metal atoms on the surface of a Ni catalyst supported on silica from the data reported by Taylor et a2.,l0 it can be seen that an increase in the Ni dispersion corresponds to an increase in activation energy. In fig. 7 the apparent activation energy, E,, for ethane hydrogenolysis is plotted as a function of In A, where A is the pre-exponential factor in the equation: (6) VE = A exp (- EJRT) PgP% where VE is the rate of ethane hydrogenolysis.With increasing activation energy a corresponding increase in In A is found. The lowest activation energy and In A values were found on the samples with the lowest dispersion. Note that the values of A used for the plot were estimated by assuming that the reaction orders do not depend upon temperature. This assumption was experimentally verified for the order of the reaction with respect to ethane. However, the order of reaction with respect to hydrogen generally assumed a smaller negative value with increasing temperature. For example, on the 4.44 wt% Ru/MgO an order of -0.9 was found at 160 OC, while a value of -0.73 was found at 189 OC; similarly on 3.0 wt % Ru/SiO, the order changed from - 2.2 at 161 OC to - 1.92 at 186 OC.For ethane hydrogenolysis such an effect has been reported A plot similar to fig. 7, showing a linear relationship between In A and E, has been found in catalytic systems exhibiting a compensation effect and various interpretations have been given to explain the phenomenon.13-16 Most of the interpretations assume l2S. GALVAGNO, J. SCHWANK, G. GUBITOSA AND G. R. TAUSZIK 2519 the existence of a heterogeneous surface with different distributions of sites Bi on which the activation energy for a particular reaction is Ei. For our Ru system, it could be suggested that the catalysts with lower dispersion and consequently higher metal coordination number mainly have sites with lower activation energy. By decreasing the metal particle size the relative amount of highly coordinated atoms decreases. The increase in the activation energy with increasing metal dispersion could be explained by assuming that the hydrogenolysis reaction on the smaller metal particles involves sites with low coordination number where the activation energy could be higher.Note also that the chemisorption of ethane or propane is dissociative, requiring the accommodation of hydrogen atoms split-off from the hydrocarbon molecules on metal sites. It is possible that the extent of dissociative removal of H atoms from the hydrocarbon differs from one catalyst to another. From the experimentally determined reaction orders, one can estimate the hydrogen content of the adsorbed hydrocarbon intermediate.2 According to this calculation the adsorbed species on the samples 4.44 wt % Ru/MgO and unsupported ruthenium seem to be less depleted of hydrogen than on the other samples. It cannot be ruled out that the less pronounced inhibiting effect of hydrogen and the lower activation energy found on 4.44 wt % Ru/MgO and on unsupported Ru is to some extent related to the smaller amount of atomic hydrogen resulting from the dissociative chemisorption of the hydrocarbon.The question has been raised in the literature12 as to whether the presence of chlorine contamination on the metal surface could lead to modifications of the hydrogenolysis activity. Generally, the use of RuCl, - H20 as a precursor in the catalyst preparation leads to significant amounts of residual chlorine in the reduced samples. Some of our samples were prepared from precursors which did not contain chlorine.Nevertheless, they show the same catalytic pattern. Therefore, it is unlikely that the large activity differences observed in our samples are due to the presence of chlorine. The results presented for the ruthenium catalysts show complex catalytic behaviour, caused by an intricate interplay of factors such as the nature of the support, the metal dispersion, the distribution of active sites, the ratio between hydrogen and hydrocarbon in the reactant mixture, the strength of adsorption and the extent of dissociative hydrocarbon adsorption. Among these factors, the metal particle size seems to play a central role in determining the catalytic characteristics.This agrees with the concept that hydrogenolysis reactions are structure-sensitive reactions. In particular, the apparent activation energies and pre-exponential factors show a clear trend, decreasing with increasing metal particle size. We will now comment further on the selectivity observed in propane hydrogenolysis. The ratio CH,/C2H6 in the reaction products was always > 1. This can be explained by the following overall scheme: Under the experimental conditions used, the rate of hydrogenolysis of C,H6 is much lower than that of propane. Furthermore, the possibility that a consecutive reaction occurs is minimized by the low conversions used. Therefore, it seems reasonable to conclude that reaction (c) does not play an important role in our system. Thus,2520 HYDROGENOLYSIS O N RU propane hydrogenolysis takes place mainly through two parallel reactions, namely (a) and (b).However as mentioned earlier, the two Ru samples with large particle size showed a much lower selectivity toward ethane formation than the other samples. For these two catalysts, an increase in temperature beyond 150 OC did not lead to a further increase in ethane formation, but resulted in a levelling off (fig. 2 and 3). Furthermore, these two catalysts showed an order of reaction with respect to hydrogen for reaction (2) close to 0 at a temperature of 154 O C . With increasing temperature, the order of reaction became positive reaching a value of ca. 0.5 at 180-190 O C (table 4). On all the other catalysts the hydrogen order remained negative.This indicates that a modified reaction mechanism contributes to the observed ' abnormal ' product distribution on ruthenium catalysts with low dispersion. The mass transfer and diffusion limitation tests allow us to rule out physical effects as being responsible for the observed phenomena. Further studies will be necessary to clarify the modified reaction mechanism. J. S. and S. G. acknowledge support for this research from the National Science Foundation (grant no. ENG 79208 180). We thank Dr Bossi for providing three of the samples, and Ms. K. Vargo, Ms. D. Radkowski, Ms. L. Flagg and Ms. L. Golze for help in the experiments. J. H . Sinfelt, Catal. Rev., 1969, 3, 175. J. H. Sinfelt, A h . Caral., 1973, 23, 91. S. Galvagno, J. Schwank, G. Parravano, F. Garbassi, A. Marzi and G. R. Tauszik, J. Catal., 1981, 69, 283. * S . Galvagno, J. Schwank and G. Parravano, J. Catal., 1980, 61, 223. J. Schwank, G. Parravano and H. L. Gruber, J. Catal., 1980,61, 19. A. Bossi, F. Garbassi, A. Orlandi, G. Petrini and L. Zanderighi, in Preparation of Catalysts IZ (Elsevier, Amsterdam, 1979), p. 405. ' I. W. Bassi, F. Garbassi, G. Vlaic, A. Marzi, G. R. Tauszik, G. Cocco, S. Galvagno and G. Parravano, J. Catal., 1980, 64, 405. P. B. Weisz, 2. Phys. Chem. (N.F.), 1957, 11, 1. J. H. Sinfelt and D. J. C. Yates, J. Catal., 1967, 8, 82. lo W. F. Taylor, J. H. Sinfelt and D. J. C. Yates, J. Phys. Chem., 1965, 69, 3857. l1 J. Barbier, A. Morales and R. Maurel, Bull. SOC. Chim. Fr., 1978, 1. l2 P. Birke, S. Engels, N. Khue and M. Wilde, 2. Chem., 1980, 20, 306. l3 A. Clark, in The Theory of Ahorption and Catalysis (Academic Press, New York, 1970), p. 260. l4 J. M. Thomas and W. J. Thomas, in Introduction to the Principles of Heterogeneous Catalysis (Academic Press, New York, 1967), p. 263. E. Cremer, Adv. Catal., 1955, 7 , 75. l6 A. K. Galwey, Ado. Catal., 1977, 26, 247. (PAPER 1/1711)

ISSN:0300-9599    

DOI:10.1039/F19827802509

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I, 1982, 78, 2521-2528 Potentiometric and Polarimetric Studies of the Reaction of Boric Acid and Tetrahydroxyborate Ion with Polyhydroxy Compounds BY J. GRAHAM DAWBER* AND DAYANG HUSNI MATUSIN Department of Chemistry and Biology, North Staffordshire Polytechnic, Stoke-on-Trent ST4 2DE Received 16th November, 198 1 Values of the association constant, K,, for the reaction of boric acid with sorbitol, mannitol, D-glucose, glycerol and ethylene glycol have been evaluated from a modified Antikainen equation using values of the dissociation constant, K*, obtained by the half-neutralisation method. This procedure gives values of K , which are easily compared with each other, and which have trends which are compatible with other previously determined values.Measurements of p P at various temperatures for the complexes of sorbitol and mannitol with boric acid allow estimates to be made of AH* and AS* for complexation. Polarimetric studies of the sorbitol and mannitol complexes in solution indicate that the tetrahydroxyborate ion, B(OH),, is more effectively complexed than boric acid and do not suggest any major differences in complexing ability between sorbitol and mannitol. The reaction of boric acid and borates with polyhydroxy compounds (polyols) has been known for many years, particularly as a means of increasing the acid strength of boric acid for its analysis by volumetric titration., The relationship between the structural characteristics of the polyol and its influence upon the acidity of boric acid was reviewed by Boiseken2T3 and it was he who suggested that the formation of boric-acid-polyol complexes occurred in two stages, first the formation of a mono- chelated borate complex, followed by a bichelated complex.It has also been suggested4-' that it is the borate ion, rather than boric acid, which is complexed by the polyol. Because the stoichiometry of the complexes is not clearly defined8-10 discrepancies occur among values derived for the association constants, since assumptions made about the complex stoichiometry lead to varying methods of calculation .5 9 1-13 In order to compare association constants it has been suggested8 that the Antikainen" equation be used, i.e. (1) where K* is the measured dissociation constant of the polyol+H,BO, solution, K , is the association constant of the complex, K , is the dissociation constant of H3BO3, C, is the stoichiometric concentration of all species containing the polyol and n is the coordination number of the complex.However, for the complexes of mannitol and glucose with boric acid8 the value of K* varied with the extent of dissociation, a mixture of acidic complexes being formed, the composition of which varied as the ratio of undissociated boric acid to borate ions changed. Similar effects are also observed with complexes of mannitol and borate To overcome this problem, Davis and Matt* used a modified Antikainen equation (in its logarithmic form) K* = Kl K , C,, + K , p(K*-KJ = -nlogl,Ci+p(KIKn) (2) 25212522 REACTIONS OF BORIC ACID WITH POLYOLS where Ci is the total stoichiometric concentration of polyol at half-neutralisation of the acid with NaOH, and K* is the dissociation constant of the polyol-boric-acid complex at half-neutralisation. pK* was obtained from the pH of the solution using the well-known Henderson-Hasselbalch equation, for which pH = pK* at half-neutralisation. The purpose of this work was to extend the studies made by Davis and Mott8 to other polyols and also to study the effect of temperature upon pK*.In addition, the formation of complexes of mannitol and sorbitol with H3BO3 and with tetrahydroxy- borate ion have been compared by polarimetric measurements. EXPERIMENTAL MATERIALS The materials used and their grades of purity were as follows: boric acid (AR), sodium chloride (AR), sodium hydroxide (AR), D-glucose (AR), glycerol (AR), ethylene glycol (GPR), mannitol (GPR) and sorbitol (GPR).The materials were used without further purification. POTENTIOMETRIC TITRATIONS Solutions of 0.05 rnol dm-3 H,BO, in 0.1 mol dm-3 NaCl were prepared containing a range of concentration of polyol from 0.01 to 0.5 mol dm-3. 50 cm3 aliquots of these solutions were titrated with 0.1 mol dmV3 NaOH. The pH titration curves were obtained using an EIL (model 7050) expanded-scale pH-meter. This model has a 2 pH-unit expanded-scale facility enabling pH readings to be estimated to 0.005, and includes both automatic and manual temperature compensation and also a fully adjustable isopotential facility. A combined glass/Ag/AgCl electrode was used to measure the pH. The values of p P were interpolated from the pH-titration curves at the half-neutralisation point.CHANGE OF pK* WITH TEMPERATURE The effect of temperature on pKy was studied for the mannitol and sorbitol complexes. Solutions of containing various amounts of polyol were prepared and standardised NaOH added to each solution corresponding to the half-neutralisation point. The pH of each solution was measured at 50, 40, 30, 25 and 20 OC, making appropriate adjustments to the temperature and isopotential controls of the pH-meter. The temperatures were maintained at - +0.5 O C over a period of ca. 2-3 min while the pH was measured. POLARIMETRIC MEASUREMENTS Solutions were made up containing 0.25 mol dm-3 mannitol plus H,BO, in amounts varying from 0 to 1 .O mol dm-3. The optical rotations of the solutions were measured with a Bellingham and Stanley model A photoelectric polarimeter using a 200 mm polarimeter tube.Angular rotations could be estimated to O.O0lo. The measurements were made at a wavelength of 435.8 nm using a low-pressure Hg lamp with other wavelengths filtered out by a cobalt glass filter plus a solution of NaN0,.14 This wavelength was used rather than the sodium D-lines since it gave larger differences in angular rotation between successive solutions. The measure- ments were repeated with the neutralised with NaOH to give B(OH),, in amounts again varying from 0 to 1.0 mol dm-3. The whole series of polarimetric measurements was then repeated using sorbitol instead of mannitol. Thus the measurements were monitoring the changes in optical rotation of the polyol, with its total stoichiometric concentration kept constant, as it complexed with increasing amounts of added H3BO3 or B(OH),.RESULTS AND DISCUSSION The pH-tritration curves for each solution were plotted and the values of pK* at half-neutralisation obtained from these graphs. The curves for sorbitol and ethyleneJ. G. DAWBER AND DAYANG HUSNI MATUSIN 2523 I 1 I I I 1 I I 5 10 15 20 25 30 35 volume of 0.1 mol dm-j NaOH/cm3. FIG. 1 .-Potentiometric titrations of 0.05 mol dm-3 H3BO3 + x mol dm-3 sorbitol where x = (a) 0, (b) 0.01, (c) 0.05, (d) 0.1, (e) 0.2, cf, 0.3, (g) 0.4 and (h) 0.5. I '*I 11 l o t I I 1 I I I I 5 10 15 20 25 30 35 volume of 0.1 mol dm-3 NaOH/cm3 FIG. 2.-Potentiometric titration of 0.05 H3B0,+x mol dm+ ethylene glycol where x = (a) 0, (b) 0.05, (c) 0.4 and (g) 0.5.0.1 (d) 0.2, (e) 0.3,2524 REACTIONS OF BORIC ACID WITH POLYOLS TABLE 1.-pP VALUES AT HALF-NEUTRALISATION FOR VARIOUS POLYOLS IN BORIC ACID SOLUTIONS p P with polyol ethylene [polyol]/mol dm-3 sorbitol mannitol D-glucose glycerol glycol 0 0.01 0.05 0.10 0.20 0.30 0.40 0.50 9.15 8.98 7.50 6.40 5.52 5.16 4.90 4.75 9.15 9.15 9.15 9.15 9.00 - 7.94 8.60 8.90 - 6.88 8.35 8.70 9.04 5.99 8.05 8.36 9.00 5.58 7.92 8.14 8.97 5.22 7.80 8.00 8.94 5.05 7.34 7.83 8-90 - - I I I 1 I I I 0.2 0.4 0.6 0.8 1.0 1.2 1.4 FIG. 3.-Antikainen plots for 0, ethylene glycol and A, sorbitol. glycol are shown in fig. 1 and 2. The p& of boric acid (pK,) was found to be 9.15, which is close to other reported l5? l6 The pK* values obtained for various polyol-boric-acid systems at half-neutralisation are given in table 1.The modified Antikainen equation [eqn (2)] predicts that graphs of p(K* -Kl) against log,, C;l should be linear; n, the coordination number, is obtained as the slope, and p(K, K,) is obtained as the intercept, from which K , can be evaluated. Linear plots were in fact obtained when the polyol concentration was greater than ca. 0.1 mol dm-3, i.e. when the polyol concentration is high relative to that of the H3BO3, and for which eqn (2) is valid. The graphs for sorbitol and ethylene glycol are shown in fig. 3. The data derved from the Antikainen plots are summarised in table 2.J. G. DAWBER AND DAYANG HUSNI MATUSIN 2525 TABLE 2.-RESULTS FROM ANTIKAINEN PLOTS [EQN (2)] slope, intercept, other values of complexant (4 PKK,) PK, Kn K?z ~ ~~ ~ sorbitol 1.95 3.80 -5.35 2.24 x lo5 - mannitol 2.27 3.99 -5.16 1.45 x lo6 1.38 x lo5" 1 x lo4 and 8.24 x 104b D-glucose 1.03 7.24 -1.91 81.3 186" 188 and 574* 8 and 77OC glycerol 1.25 7.32 - 1.83 67.6 36.4 and 81.3b 16.0 and 41.2c ethylene glycol 0.67 8.98 -0.17 1.48 1.85 and 0.lc a Ref.(8); ref. (11); ref. (5). TABLE 3.-pK* FOR HaBOa + SORBITOL AT VARIOUS TEMPERATURES pK* at [sorbitol]/mol dm+ T/OC 0 0.05 0.1 0.2 0.5 50 8.86 7.50 6.54 5.76 4.98 40 8.94 7.44 6.44 5.65 4.85 30 9.04 7.35 6.28 5.50 4.70 25 9.09 7.28 6.18 5.38 4.58 20 9.15 The values of n when glucose, glycerol and ethylene glycol were used as complexants correspond to the formation of a 1 : 1 complex with boric acid. For sorbitol and mannitol, however, n is ca.2, indicating the formation of a bichelated complex. The values of the association constants (K,) for the various complexants, evaluated from the Antikainen equation, are compared in table 2. There is broad agreement between the values of K , evaluated by this method and those determined by other methods for which certain association equilibria are assumed. However, the Antikainen method, based upon the p P at half-neutralisation,8 allows a comparison to be made between various complexants without assumptions having to be made concerning the association equilibria. The value of K , for sorbitol is only slightly larger than that for mannitol, indicating that there is not a major difference in the complexing abilities of these two polyols for boric acid.The values of K , for the other polyols studied (table 2) are considerably lower than those for sorbitol and mannitol and follow trends similar to those found in other st~dies.~, 8 v l1 The values of p P at various temperatures and at various concentrations of sorbitol and mannitol are given in tables 3 and 4. It can easily be shown from the van't Hoff isochore that p P = AH*/2.303 RT+constant. Thus the enthalpy of ionisation of the polyol-H,BO, complex, AH*, can be obtained from a graph of pK* against 1/T. The values of AH* obtained in this way are plotted2526 REACTIONS OF BORIC ACID WITH POLYOLS TABLE 4.-pK* FOR H3B03 +- MANNITOL AT VARIOUS TEMPERATURES p P at [mannitol]/mol dm-3 T/OC 0 0.05 0.1 0.2 0.5 50 8.86 7.92 7.08 6.20 5.30 40 8.94 7.86 7.00 6.10 5.20 30 9.04 7.84 6.83 5.96 5.08 25 9.09 7.82 6.75 5.87 4.98 I,,,,,,,,,,,,, 2 4 6 8 1 0 1 2 FIG.4.-Enthalpy of ionisation of H,BO, complexes with 0, sorbitol and x , mannitol. mJmt as a function of polyol concentration for mannitol and sorbitol in fig. 4. Although the AH* values will not be very precise, they do show the trends occurring as complexation takes place. The values of AH* level out at a ratio of ca. two moles of complexant per mole of H,BO,, i.e. rn,/rn, x 2, which is consistent with the formation of a bichelated complex, with AH* for sorbitol being slightly more exothermic than for mannitol. Assuming that AGe* = -RTln P, it is possible to calculate the corresponding values of AS,*,, for formation and ionisation of the complex, and these values for sorbitol and mannitol are given in table 5.The As&, values will only be very approximate, but nevertheless the large negative values do indicate the large loss of vibrational and rotational freedom of motion of those groups in the polyol which become involved in the complexation with the boron atom.J. G. DAWBER AND DAYANG HUSNI MATUSIN 2527 TABLE 5.-VALUES OF AS* FOR SORBITOL AND MANNITOL COMPLEXES AS,*,,/J K-l mol-l rnz/mla sorbitol mannitol 0 -115 -115 1 -195 -173 2 -210 -206 4 -197 -192 10 -183 -172 a Moles of polyol per mole of H,BO,. TABLE 6.--EFFECT OF BORIC ACID (a) AND BORATE ION (b) ON OPTICAL ROTATION OF SORBITOL (mz = 0.25 mol dm-3) AT 436 nm a/lO-, O m2 mol-I mla/mol dm-, ml/m2 (a) (4 0 0.02 0.04 0.10 0.20 0.30 0.40 0.50 0.60 0.70 0.80 0.90 0.96 1 .oo 0 0.08 0.16 0.40 0.80 1.20 1.60 2.0 2.4 2.8 3.2 3.6 3.84 4.0 - 5.75 - 5.50 - 5.04 - 4.00 - 2.52 - 1.26 - 0.02 + 0.82 1.76 2.66 3.44 3.94 4.84 - - 5.75 - 1.20 + 6.64 15.70 24.98 30.18 33.1 35.6 37.3 38.3 39.8 40.1 40.2 - a Moles of (a) boric acid and (b) borate ion.The optical rotation results were converted to molar optical rotation, a!,, by a, = a/cl, where a is the measured rotation, 1 the path-length of the polarimeter tube (in m) and c the total polyol concentration which for these measurements was 250 mol m-3, i.e. 0.25 mol dm-3. The units of a, are thus O m2 mol-l. The experiments were conducted with the polyol concentration (m,) kept constant and with increasing amounts of boric acid or tetrahydroxyborate ion, m,, added up to a maximum concentration of 1 mol dm-3 (i.e.ml/m2 = 4 for the most concentrated solutions). The optical rotation results are compared in tables 6 and 7. The results suggest that the B(0H); ion, rather than boric acid, complexes with the polyol. The configuration of B(0H); will be tetrahedral and this will involve rather less change in symmetry on complexing with the polyol compared with H3BO3, for which the bond angles are likely to be ca. 120°. The data show that the formation of the mannitol-borate complex is accompanied by larger changes in optical rotation than in the case of the sorbitol-borate complex, which suggests the mannitol complex as having the greater asymmetry. The2528 REACTIONS OF BORIC ACID WITH POLYOLS TABLE ?'.-EFFECT OF BORIC ACID (a) AND BORATE ION (b) ON ROTATION OF MANNITOL (m, = 0.25 mol dm-3) AT 436 nm ~c/lO-~ O m2 mol-l mla/mol dm-3 ml/m2 (a) 0 0.04 0.10 0.20 0.30 0.40 0.50 0.60 0.70 0.80 0.90 1-00 0 0.16 0.40 0.80 1.20 1.60 2.0 2.4 2.8 3.2 3.6 4.0 -2.1 10 - 1.080 + 0.860 3.740 6.100 8.72 10.96 13.34 15.44 17.28 19.32 21.76 - 2.14 +7.10 24.02 53.9 80.1 93.4 100.7 106.7 1 1 1.0 113.6 117.2 120.0 a Moles of (a) boric acid and (b) borate ion.changes in optical rotation as a function of mJm, do not show any major differences between sorbitol and mannitol in the effectiveness of complexation with H3BO3 and B(0H);. The helpful comments of a referee are gratefully acknowledged. M. G. Mellon and V. N. Morris, Ind. Eng. Chem., 1924, 16, 123. J. Boeseken, Adu. Carbohydr. Chem., 1949, 4, 189; 1949, 12, 81. J. P. Sickels and H. P. Schultz, J. Chem. Educ., 1964, 41, 343. J. Boeseken and N. Vermaes, J. Phys. Chem., 1931,35, 1477. A. Deutsch and S . Asoling, J. Am. Chem. SOC., 1949, 71, 1637. G. L. Roy, A. L. Lafferiere and J. 0. Edwards, J. Znorg. Nucl. Chem., 1957, 4, 106. S. D. Ross and A. J. Catotti, J. Am. Chem. SOC., 1949, 71, 3563. T. P. Onak, H. L. Landesman, R. E. Williams and I. Shapiro, J. Phys. Chem., 1959, 63, 1533. * H. B. Davis and C. J. B. Mott, J. Chem. SOC., Faraday Trans. 1, 1980, 76, 1991. @ R. Larsson and G. Nunziata, Acta Chem. Scand., 1970, 24, 2145. lo M. Mazurek and A. S. Perlin, Can. J. %hem., 1963, 41, 2403. l1 P. Antikainen, Ann. Acad. Sci. Fenn., Sect. AZZ, 1954, 56, 3; Suomen Kemistil. B, 1958, 31, 255. l2 S. J. Angyal and D. J. McHugh, J. Chem. SOC., 1957, 1423. l3 J. J. Kankare, Anal. Chem., 1973, 45, 2050. l4 J. G. Dawber, J. Chem. SOC., Faraday Trans. I , 1978, 74, 960. l5 A. Albert and E. P. Serjeant, The Determination oflonisation Constants (Chapman and Hall, London, l6 Handbook ofChemistry and Physics, ed. R. C . Weast (C.R.C. Press, Columbus, Ohio, 60th edn, 1980). 1971). (PAPER 1 / 1776)

ISSN:0300-9599    

DOI:10.1039/F19827802521

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1982, 78, 2529-2534 Viscosities of Gaseous Argon-Nitrogen Mixtures BY G. PETER MATTHEWS, HAIDEE SCHOFIELD, E. BRIAN SMITH* AND ANDREW R. TINDELL Physical Chemistry Laboratory, South Parks Road, Oxford OX1 3QZ Received 20th November, 198 I The viscosities of gaseous argon-nitrogen mixtures have been measured in the temperature range 120-1600 K using a capillary flow technique. The results are considered to be accurate to f 1 % and are used to calculate interaction viscosity coefficients and binary diffusion coefficients for the mixtures. A knowledge of the transport properties of gases is of considerable value in the study of intermolecular forces. The recent development of methods of inverting gas transport-property measurements and gas imperfection has played a significant role in advancing our knowledge of monatomic species5-’ and has complemented information obtained from spectroscopic8 and molecular-beam studie~.~-ll The extension of inversion techniques to anisotropic interactions is a problem under active investigation at the present time.l2?l3 In order to provide data suitable for these studies, accurate measurements of the transport properties of simple anisotropic systems are required. In this paper we describe the .determination of the viscosities of argon-nitrogen mixtures over a very wide range of temperature (1 20- 1600 K).EXPERIMENTAL The viscosities of mixtures of argon and nitrogen were measured using two capillary flow viscometers that have been described fully in previous p~b1ications.l~. l5 One viscometer was used to determine viscosities from liquid-nitrogen temperatures to just above room temperature.The other was used for the temperature range 300-1600 K. Both sets of apparatus were of similar design, comprising two thermostatted glass vessels connected to one another by a coiled capillary tube. In the high-temperature apparatus the capillary tube is enclosed in a furnace, and in the low-temperature version it is contained in a cryostat in order to maintain constant temperature. The times taken for gas to pass from one vessel to the other, with a known pressure difference across the capillary tube, were measured. These times of flow were monitored by a mercury manometer fitted with electrical contacts. The time for the mercury level to drop past successive contacts was recorded on a series of automatic timers.On the high-temperature apparatus both silica and platinum-rhodium capillary tubes, operating over different temperature ranges, were employed to minimise the various corrections for non-Poiseullian flow. Only a single glass capillary tube proved necessary in the low-temperature apparatus. The argon-nitrogen mixtures used in the experiments were obtained from two sources. In most of the hgh-temperature experiments the pure gases supplied by the British Oxygen Co. were mixed in a 10 dm3 globe and the compositions were determined by means of an accurate gas density balance and confirmed by mass spectrometry. In the majority of the low-temperature experiments the argon-nitrogen mixtures were supplied directly by the British Oxygen Co.The composition of the mixtures was determined by B.O.C. (Special Gases Division) and on receipt was checked using a mass spectrometer. Analysis of the pure components showed the argon to be > 99.9% pure (major impurity nitrogen), and the nitrogen to be > 99.99% pure. Several crucial though small corrections must be made to the times of flow to allow for curved 82 2529 FAR 12530 VISCOSITIES OF Ar-N, MIXTURES pipe flow, kinetic energy effects, gas imperfections and non-zero gas flow velocity at the walls of the tube. After all these corrections have been made it is possible to obtain the ratio of the argon-nitrogen mixture to that of a standard gas at the same temperature. In this work argon was employed as the reference gas. The values used are given in table 1 and can be expressed by the relation lnq=AInT+B/T+C/T2+D (1) where the coefficients for both high and low temperatures are given in table 2.room temperature and Gough et d . 1 6 at low temperatures. All are based on the value This standard is consistent with the proposed nitrogen standards of Dawe and Smith15 above qN2 (293.2 K) = 175.7 x lo-' kg m-l s-l. RESULTS Experiments were carried out at eleven temperatures, five on the high-temperature viscometer in the range 400-1600 K and six using the low-temperature viscometer. For both the high- and low-temperature experiments mixtures of three compositions were employed with nominal mole fractions x of 0.25, 0.50 and 0.75. The results for these mixtures and the pure components are given in table 1.In the study of dilute gas mixtures it is convenient to interpret the results in terms of the so-called interaction viscosity coefficient qI2 which refers specifically to the unlike interactions in the mixture. Interaction viscosities may be calculated17 from a knowledge of the viscosities of the mixture and the pure components at the same temperature, together with a dimensionless quantity A:2 which is the ratio of two collision integrals Values of A:2 are relatively insensitive to the potential function used to calculate them, and we have employed the BBMS potential5 (which gives an accurate representation of the intermolecular forces of argon) to calculate the values given in table 1 . Interaction viscosities were calculated for mixtures of all three compositions at each temperature.The values were not completely independent of composition at a given temperature owing to uncertainties in the measured viscosities and mole fractions and deficiences in the first-order kinetic theory expression used in the calculations. The average values of qI2 at each temperature are also given in table 1. COMPARISON WITH OTHER WORK The only measurements recorded in the literature are by Kestin and co-workers.18- l9 They recorded mixture viscosities for four compositions over the temperature range 293-767 K. Their data were analysed to give the interaction viscosities as described above. The resulting interaction viscosities from the work of Kestin together with the results of the present work were fitted to a smooth curve of the form of eqn (1) over the two temperature ranges 120-300 K and 300-1600 K.The coefficients for this curve are shown in table 2 and the interaction viscosities calculated from the various sources are shown in table 3. The percentage deviations of the viscosities from the smooth curve from the data of table 3 were calculated and the deviation plot is shown in fig. 1. The r.m.s. percentage deviation of all the points from the curve is O S % , with the largest deviation of any single point being ca. 1%. These deviations are well within the expected error limits given the uncertainties involved in the calculation of the interaction viscosities. The results constitute a consistent set of results over the temperature range 120- 1600 K.G.P. MATTHEWS, H. SCHOFIELD, E. B. SMITH AND A. R. TINDELL 2531 TABLE 1 .-VISCOSITIES OF Ar-N, MIXTURES 118.90 - - 151.14 - - 203.94 - - 240.18 293.2 - - - 298.18 - - 400.0 - - 600.0 - - 797.9 - - - 1200.7 - - 1597.4 - - 1.1 110 - - 1.1070 - - 1.1034 - - 1.1030 1.1034 - - - 1.1043 - - 1.107 - - 1.117 - - 1.126 - - - 1.140 - - 1.151 - - 96.9 - - 122.4 - - 162.3 - - 188.0 223.3 - - - 226.5 - - 288.3 - - 389.8 - - 475.3 - - - 623.8 - - 749.8 - - 80.2 - - 100.6 - - 131.1 - - 150.2 176.3 - - - 178.1 - - 222.9 - - 296.0 - - 357.5 - - - 467.5 - - 556.1 - - 0.2490 0.4970 0.7506 0.2490 0.4970 0.7506 0.2490 0.4970 0.7506 0.2490 0.4970 0.7441 0.4973 0.2506 0.2490 0.4970 0.7506 0.7441 0.4973 0.2506 0.7441 0.5094 0.2506 0.5094 0.7246 0.5147 0.2747 0.7355 0.5203 0.2130 0.7547 0.5067 0.2562 93.4 89.4 85.2 118.2 113.6 107.9 154.8 148.6 140.2 179.2 170.4 189.5 202.1 213.3 2 16.2 203.5 191.7 241.3 259.2 275.0 322.8 346.6 368.5 419.1 388.5 4 14.4 447.3 509.0 539.9 589.8 606.5 658.9 704.3 88.4 88.0 88.8 112.4 113.3 114.1 144.6 147.1 147.2 167.5 168.4 198.7 200.0 200.4 202.9 200.8 202.3 254.7 257.1 259.9 343.2 344.9 341.5 41 5.3 404.2 407.4 418.5 534.6 529.1 532.1 647.0 653.5 649.7 88.4 - - 113.3 - - 146.3 - - 167.9 199.7 - - - 202.0 - - 257.2 - - 343.2 - - 41 1.3 - - - 53 1.9 - - 650.1 - - TABLE 2.-cOEFFICIENTS OF THE CURVE In Fj‘ F A In T+ B/ T+ C/ T2 + D FOR THE ARGON STANDARD AND FOR ARGON-NITROGEN INTERACTION VISCOSITTES TIK A B c D Ar 120-300 0.583 152 -96.1924 2 923.99 2.38966 300- 1700 0.599 369 - 57.5041 - 3 118.53 2.236 300-1600 0.707 301 101.2348 -23 859.5 - 1.201 747 Ar-N, 120-300 1.380 227 237.3904 - 12 696.7 -3.21 1 58 82-22532 VISCOSITIES OF Ar-N, MIXTURES TABLE 3.-INTERACTION VISCOSITIES FOR ARGON-NITROGEN MIXTURES ?ld lo-' T / K source kg m-l s-l 118.9 151.14 203.94 240.18 293.00 292.18 298.00 298.18 303.00 367.00 400.00 467.00 57 1 .OO 600.00 673.00 767.00 797.9 1200.7 1597.4 this work this work this work this work ref.(19) this work ref. (18) this work ref. (19) ref. (18) this work ref. (18) ref. (18) this work ref. (18) ref. (18) this work this work this work 88.40 113.30 146.30 167.90 197.73 199.70 200.52 202.00 207.40 237.39 257.20 285.21 327.32 343.20 366.15 400.69 41 1.30 53 1.90 650.10 1.0 l . 7 0 0 0 - 2 . 0 ~ 200 600 1000 1400 TIK FIG. 1 .-Deviations in interactive viscosities (qI2) for argon-nitrogen mixtures.0, This work; +, ref. (18); V, ref. (19).G. P. MATTHEWS, H. SCHOFIELD, E. B. SMITH AND A. R. TINDELL 2533 DISCUSSION Binary gaseous diffusion coefficients may, to first order, be directly calculated from A&RT PP interaction viscosities by2* [Dl211 = % t t l Z l l _ _ _ where the subscripts indicate that the equation involves only first-order results for D,, and q,,, p is the reduced mass and A,*, the ratio of collision integrals defined earlier. This relation may be used to calculate diffusion coefficients from viscosity data and, when experimental diffusion data are available, to check on the consistency of diffusion and viscosity measurements. Storvick and Mason2' have shown that eqn (2) gives values of D,, that correspond most closely to the limit when the heavier component is present only as a trace.Mason and Marrero,, have shown that the second approximation for D,, can be represented by where A,, is a correction which involves the concentration dependence of Dl,. Mason and Marrero give a semi-empirical expression for the factor where Cy, = 52*(1,2)/52*(1,1), x, is the mole fraction of component 1 and a, b and 6 are assigned the values 0.029, 0.10 and 1.0, respectively. The results of these calculations are given in table 4 together with recommended values of D,, in the TABLE 4.-RESULTS FOR ARGON-NITROGEN BINARY DIFFUSION COEFFICIENTS 293.2 1.1035 0.8947 0.193 298.18 1.1026 0.8957 0.199 400.0 1.1078 0.9100 0.341 600.0 1.1 178 0.923 1 0.688 797.9 1.1268 0.9279 1.077 1 200.7 1.1411 0.9300 2.178 1597.4 1.1519 0.9290 3.575 r.m.s.deviation from expt (%): 2.89 0.194 0.190 0.199 0.196 0.342 0.327 0.690 0.666 1.082 1.097 2.188 2.246 3.593 3.703 2.88 a Calculated for an equimolar mixture; from ref. (23). temperature range 244-10000 K given by Marrero and Mason.23 The experimental uncertainties in D12 are quite large, being in the range from 2% at 300 K to 7% at 1000 K. The two sets of results are therefore consistent given the uncertainties involved. The principal purpose of the measurements described above is to provide data which can be used to elucidate further the argon-nitrogen potential-energy function. This elucidation is the subject of an investigation currently under way.2534 VISCOSITIES OF Ar-N, MIXTURES ' G.C. Maitland, M. Rigby, E. B. Smith and W. A. Wakeham, Intermolecular Forces - Their Origin and Determination (Oxford University Press, Oxford, 1981), chap. 9. D. W. Gough, G. C. Maitland and E. B. Smith, Mol. Phys., 1972, 24, 151. H. E. Cox, F. W. Crawford, E. B. Smith and A. R. Tindell, Mol. Phys., 1980, 40, 705. E. B. Smith, A. R. Tindell, B. H. Wells and F. W. Crawford, Mol. Phys., 1981, 42, 937. G. C. Maitland and E. B. Smith, Mol. Phys., 1971, 22, 861. D. W. Gough, E. B. Smith and G. C. Maitland, Mol. Phys., 1973,25, 1433. D. W. Gough, E. B. Smith and G. C. Maitland, Mol. Phys., 1974,27, 867. D. E. Freeman, K. Yoshino and Y. Tanaka, J. Chem. Phys., 1973,59, 5160, and references therein. J. M. Parson, P. E. Siska and Y. T. Lee, J. Chem. Phys., 1972, 56, 151 1, and references therein. lo C. H. Chen, P. E. Siska and Y. T. Lee, J. Chem. Phys., 1973, 59, 601. l1 K. M. Smith, A. M. Rulis, G. Scoles, R. A. Aziz and V. Nain, J. Chem. Phys., 1977, 67, 152. I2 E. B. Smith, D. J. Tildesley, A. R. Tindell and S . L. Price, Chem. Phys. Lett., 1980, 74, 193. l3 E. B. Smith and A. R. Tindell, Faraday Discuss. Chem. SOC., 1982, 73, in press. I4 A. G. Clarke and E. B. Smith, J. Chem. Phys., 1968,48, 3988. I6 R. A. Dawe and E. B. Smith, J. Chem. Phys., 1969, 52, 693. Is D. W. Gough, G. P. Matthews and E. B. Smith, J , Chem. Soc., Faraday Trans. I , 1976, 72, 645. l7 J. 0. Hirschfelder, C. F. Curtiss and R. B. Bird, Molecular Theory of Gases and Liquids (Wiley, New York, 1954), p. 529. J. H. Hellemans, J. Kestin and S. T. Ro, J. Chem. Phys., 1972, 57, 4038. l* R. DiPippo, J. Kestin and K. Oguchi, J. Chem. Phys., 1967, 46, 4758. 2o J. 0. Hirschfelder, C. F. Curtiss and R. B. Bird, Molecular Theory of Gases and Liquids (Wiley, New 21 F. Storvick and E. A. Mason, J. Chem. Phys., 1966, 45, 3752. 22 E. A. Mason and T. R. Marrero, Adv. At. Mol. Phys., 1970, 6, 156. 23 T. R. Marrero and E. A. Mason, J. Phys. Chem. Ref. Data, 1972, 1, 3. York, 1954), p. 540. (PAPER 1 / 1805)

ISSN:0300-9599    

DOI:10.1039/F19827802529

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I, 1982, 78, 2535-2546 Reactions of Oxygenated Radicals in the Gas Phase Part 9.-Self-reactions of Isopropylperoxy Radicals BY LESLIE T. COWLEY Shell Research Limited, Thornton Research Centre, P.O. Box 1, Chester CH1 3SH AND DAVID J. WADDINGTON* AND ALLAN WOOLLEY~ Department of Chemistry, University of York, Heslington, York YO1 5DD Received 23rd November, 198 1 The principal products of the photo-oxidation of trans-2,2'-azopropane between 333 and 373 K are acetone, isopropyl alcohol, isopropyl hydroperoxide, acetaldehyde, formaldehyde, methyl alcohol and cis-2,2'-azopropane. The reaction mechanism has been simulated in detail, and, in conjunction with results obtained earlier for the overall self-reaction of isopropylperoxy radicals, the following rate data have been obtained for the reactions (3 4 (3 b) 2(CH,),CHO, - --+ (CH,),CHOH + (CH,),CO + 0, 2(CH,),CHO; -+ 2(CH,),CHO* +O, k3Jk3, increases with temperature, from 1.39kO.04 at 302 K to 1.83 f0.04 at 333 K and 2.80f0.08 at 373 K.Values of A,, and ABb of 2.44f0.31 x lo7 and 1.381t0.26 x lo9 dm3 mol-' s-l and E,, and E3b of 12.05 1.0 and 21.3+ 1.5 kJ mo1-l were determined. In this series of papers1q2 the importance of obtaining reliable rate data for the reactions of oxygenated radicals in the gas phase, so that reaction mechanisms for combustion and atmospheric processes may be rigorously tested, has been em p hasised . In two of these papers rate data have been obtained for the overall effective bim olecular sel f-reac tions of isoprop ylper ox y radicals, using u .v.spectroscopic techniques,' over a temperature range 302-373 K, 2(CH3),CH0, -+ products ( o w and for the individual processes2 2(CH3),CH0, -+ (CH,),CHOH + (CH,),CO + 0, 2(CH3),CHO; --+ 2(CH3),CH0 + 0, (3 4 (3 4 at 302 K. This provided, for the first time, values of elementary rate constants, directly determined, for these two reactions. The shortage of elementary data is seen to be particularly unfortunate when it is remembered that measurements of overall self-reaction rate constants have raised as many fascinating problems as they have solved. For example, the value of the overall rate constant depends on the nature of the alkyl group of the alkylperoxy radical involved, for which present theory is unable to account successfully. The problem may be more tractable if the individual, rather than the overall, rate parameters are considered.Present address: University College School, Frognal, London NW3 6XH. 25352536 SELF-REACTIONS OF ISOPROPYLPEROXY RADICALS Further, once the Arrhenius parameters of these individual reactions are known, it should be possible to obtain rate data for other reactions of alkylperoxy radicals, e.g. abstraction and addition reactions, using the self-reactions as the control in competition with the reaction being studied. In this way, important data may be obtained to help us understand the complex behaviour of alkylperoxy radicals in gas-phase reactions. In this paper results for the self-reactions of isopropylperoxy radicals over the temperature range 302-373 K are presented.EXPERIMENTAL The apparatus and methods of analysis of products and reactants were described in an earlier paper.* As before, trans-2,2’-azopropane was prepared from isopropylamine and purified to > 99% (by gas chromatography); no cis-2,2’-azopropane was dete~table.~ RESULTS The products of photolysis of tran~-2,2’-azopropane, both in the presence and absence of oxygen, were confirmed by gas chromatography and mass-spectral analyses. In the absence of oxygen, the products at 333 and 373 K are similar to those found at 302 K,2 &-2,2’-azopropane, propane, propene and 2,3-dimethylbutane (fig. 1). The 2.0 4.0 6.0 8.0 10.0 t/103 FIG. I .-Photolysis of trans-2,2’-azopropane. 2,2’-tran.s-azopropane, 5 Torr; nitrogen, 495 Torr. 0, m, 2,3-Dimethylbutane at 333 and 373K; 0, @, propene and propane at 333 and 373 K ; A, A, ci,s-2,2’-azopropane at 333 and 373 K.( 1 Torr = 101 325/760 Pa.) rates of formation of propane and propene are indistinguishable and the ratio of the rates of formation of these products to that of 2,3-dimethylbutane is constant as reaction proceeds, but varies slightly with temperature (fig. 1). The rate of formation of cis-2,2’azopropane decreases with time until a stationary concentration of the isomer is achieved. Some of the photolysis products in the presence of oxygen are similar to those at 302 K , ci.~-2,2’-azopropane, acetone, isopropyl alcohol and isopropyl hydroperoxide.2.0 4.0 6.0 8.0 10.0 ti103 s 2.0 4.0 6.0 8.0 10.0 t1103s 2.0 4.0 6.0 8.0 10.0 t/103s 2.0 4-0 6.0 8.0 100 ti103 FIG.2.-Photo-oxidation of 2,2’-azopropane at 333 K. 0, Acetone; 17, isopropyl , hydroperoxide; 0, isopropyl alcohol; 0 , acetaldehyde; 0, formaldehyde; 0, methyl alcohol. Experimental points are shown; the lines are from calculations following simulations (a) 2,2’-Azopropane, 5 Torr; oxygen, 10 Torr; nitrogen, 485 Torr. (b) 2,2’-Azopropane, 5 Torr; oxygen, 50 Torr; nitrogen, 445 Torr. (c) 2,2’-Azopropane, 5 Torr; oxygen, 495 Torr. (6) 2,2’-Azopropane, 5 Torr; oxygen, 10 Torr; nitrogen, 485 Torr.2538 SELF-REACTIONS OF ISOPROPYLPEROXY RADICALS 3*0r 2.0 4.0 6.0 8.0 10.0 t i 1 0 3 s r I 2.0 4.0 60 8.0 10.0 12.0 t 1 1 0 3 S 2.0 4.0 6.0 8.0 10.0 12.0 t1103s FIG. 3.-Photo-oxidation of 2,2’-azopropane at 373 K. (a) 2,2’-Azopropane, 5 Torr; oxygen, 25 Torr; nitrogen, 470 Torr.(b) 2,2’-Azopropane, 5 Torr; oxygen, 200 Torr; nitrogen, 295 Torr. (c) 2,2’-Azopropane, 5 Tom; oxygen, 495 Torr. For key, see fig. 2.L. T. COWLEY, D. J . WADDINGTON AND A. WOOLLEY 2539 However, at 333 and 373 K significant amounts of formaldehyde, acetaldehyde and methyl alcohol are also formed. Propane and 2,3-dimethylbutane were not observed; however, traces of propene were detected. The rates of formation of all the oxygenated products vary with photolysis time and with temperature (fig. 2 and 3). The initial oxygen pressure was varied, keeping the total pressure (with nitrogen) constant. As the oxygen pressure is increased, the yield of isopropyl alcohol decreases while that of isopropyl hydroperoxide increases. The sums of their concentrations equals that of acetone at high oxygen pressures, the formation of formaldehyde, acetaldehyde and methyl alcohol being suppressed on increasing the oxygen pressures.As at 302 K,, the rate of decomposition of 2,2’-azopropane was determined by measuring the rates of formation of hydrocarbons, at constant total pressure. This gave a rate constant for the reaction of 1.5 x s-l, which was independent of temperature. DISCUSSION The initiation reaction for the photo-oxidation of trans-2,2’-azopropane is c~mplex.~-~ However, we are only concerned with the subsequent reactions of isopropyl radicals, and the complexity of the initiation processes do not affect directly the kinetics of these reactions., On photolysis of 2,2’-azopropane, as at 302 K, equimolar proportions of propene and propane are formed, the yields of the C, hydrocarbons being smaller than that of 2,3-dimethylbutane.The three products are formed from self-reactions of isopropyl radicals formed from reaction (1): (CH,),CHN,CH(CH,), -+ 2(CH,),t]H + N,. (1) The ratio of rate constants for the disproportionation and combination reactions, k,/k,, alters slightly on increasing the temperature. Our value of 0.45 & 0.20 exp (- 676 & 120/RT) compares with values obtained earlier of 0.4+ 0.1 exp (- 1090 f 105/RT)s and 0.36 exp (+ 1 130/RT),9 where T is in K. Our figures are found for less extensive data than those used by earlier workers; they serve as a check on our analytical procedures and as a means of obtaining a value of k,. The results for the photo-oxidation of trans-2,2’-azopropane correspond to the following mechanism :, RN,R -+ 2R* +N, R*+O,(+M)+ RO,*(+M) 2R0, + (CH,),CO + (CH,),CHOH + 0, 2RO; + 2RO* +O, (3 a) (3 b) 2RO* -+ (CH,),CO+(CH,),CHOH ROO +O, -+ (CH,),CO+HO,* RO,.+HO; -+ RO,H+O, (6) RO; + ROO -+ (CH,),CO + RO,H (7) ROO + R0,H + (CH,),CHOH + RO; (8)2540 SELF-REACTIONS OF ISOPROPYLPEROXY RADICALS ROO + HO, -+ (CH,),CHOH + 0, (9) HO; +HO; -+ H,O,+O,* (10) (where R = isopropyl). A computer model of the system,,? lo designed to simulate the time-composition behaviour of a homogeneous gas-phase reaction, was used in which a numerical integration procedure was employed for the solution of ' still ' differential equations. An optimisation procedure was used to find the ratio k,/k,. This model allowed for the effect of initial oxygen pressure on the relative yields of products, and from it a value of k3,/k3, was obtained of 1.39 at 302 K., As kobs had been determined directly spectroscopically it was possible to obtain values of k,, and k,, independently., The most striking difference in product distribution between the results obtained at 302 K and in the temperature range reported in this paper, 333-373 K, is the formation of significant amounts of formaldehyde, acetaldehyde and methyl alcohol.The formation of C, and C, compounds requires that a carbon-carbon bond is broken, and the logical precursor is the isopropoxy radi~a1.ll-l~ On decomposition, acetaldehyde and a methyl radical are formed: (1 1) ~H,+O,(+M)+CH,O; (+M). (12) (CH,),CHO (+ M) -+ CH,CHO + CH, (+ M).Methyl radicals react rapidly with oxygen to yield methylperoxy radicals : The introduction of methylperoxy radicals into the system complicates the reaction mechanism, for they may undergo all the reactions suggested for isopropylperoxy radicals and, further, the two forms of alkylperoxy radicals will undergo cross-reactions. Thus one has to add reactions (13a) and (13b)-(21) to the scheme: 2Me0,. -+ 2Me0 + 0, 2Me0; + HCHO + MeOH + 0, MeO,*+RO,* -+ MeO* +RO* +O, MeO; + RO; -+ HCHO + ROH + 0, MeO, + RO, -+ (CH,),CO + MeOH + 0, (134 (13@ (14d (144 (144 RO, + MeO* -+ HCHO+ R0,H Me00 +O, + HCHO+HO; MeO; +ROO + (CH,),CO+MeO,H (17) MeO* + RO,H -+ MeOH + RO; (18) ROO +MeO,H -+ ROH+MeO,* MeO* + Me0,H -+ MeOH + MeO, MeO; + HO,* -+ Me0,H + 0,. (21) There is no shortage of rate constants quoted for some of these reactions, and one has to exercise considerable care in choice when using them in simulation procedures. It is important not only to be able to justify choice but also to test them in the mechanism.In this work we tested the mechanism by varying reaction pressures, total pressure and temperature; in addition there were, as will be described later, some further experiments.L. T. COWLEY, D. J. WADDINGTON AND A. WOOLLEY 2541 Experience of simulations of the system at 302 K indicated that mutual and cross-reactions involving only alkoxy and hydroperoxy radicals are of no importance to the overall kinetics of the system. Initial simulations suggested that this is also so at 333 and 373 K. Indeed, the predicted concentrations of methoxy radicals are even lower than those of either isopropoxy or hydroperoxy radicals.Bimolecular reactions involving only isopropoxy, methoxy and hydroperoxy radicals were, therefore, omitted from the present mechanism and subsequent computer simulations. The predicted concentrations of methylperoxy radicals are also very low, typically 3 x that of isopropylperoxy. The reaction between these radicals and others in very low concentration (alkoxy and hydroperoxy) are therefore also of no significance during simulations, and were also omitted from the mechanism. Nevertheless, in order to test these assumptions in simulations, rate constants for the reactions had to be chosen. Values chosen for the rate constants for reactions (2), (4), (6), (7), (9) and (10) are discussed in an earlier paper.2 Reactions (15), (17) and (21) are exothermic radical- radical reactions, and values of between 10, and lo9 dm3 mol-' s-, are reasonable.It was found that by setting k, to be 9.2 x 10, dm3 mol-1 s-l the absolute magnitudes of k, and k , are determined; these in turn give good agreement between the simulation and experimental results over all oxygen pressures studied. The precise values of k,,, k,, and k,, do not affect the overall simulation results and were set equal to k,. For results at temperatures above 302 K, the key reaction to add to the mechanism is reaction (1 1): (CH,),CHO=( + M) + CH,CHO + cH3( + M). (1 1) There is a wide range of quoted values for the Arrhenius parameters for this reaction.l39 l5 There is even confusion on whether the A factor increases or decreases on increasing the pressure of the system.The most satisfactory values for k,,, from our simulations, were 1.39 x 10, s-l at 333 K and 1.20 x lo3 s-l at 373 K, at the pressures used. These values are in broad agreement with those found later in separate experiments in which the effect of changing pressure on reaction (1 1) was studied in detail.16 Values used for k13, and k13b are those determined by Parkes.17 These values were obtained at room temperature, but Parkes18 suggests that the activation energy for these reactions is close to zero. Although Parkes suggests that there is a further terminating reaction in which methyl peroxide is formed, the compound was not detected under any of the conditions studied.As the experimental results also tallied with product studies of photo-oxidation of azoethane,lg it seems that the contribution from this reaction is small, and it has been left out of the present simulations. Parkesl, suggests that the rate constant for the cross-reaction between methylperoxy and t-butylperoxy radicals is closer to the value of that for the self-reactions of methylperoxy radicals than of that for the self-reaction of t-butylperoxy radicals. Thus values used for k14,, k14, and k14c have been arbitrarily weighted by a factor of 2 towards the corresponding self-reaction of methylperoxy radicals. In the absence of further evidence this procedure seems as satisfactory as any other. An error of much less than an order of magnitude in the values used for these rate constants is suggested.Such an error will not have a profound effect upon the simulations as the alternative self-reactions lead to the same products. Absolute rate constants for reactions ( 5 ) and (8) at 302 K, and at higher temperature reactions (18), (19) and (20), are difficult to assign. However, as a first step, an optimisation procedure was used2 to obtain relatively precise values of the ratio k,/k, of 166 f 5, 103 f 5 and 62 f 3 at 302, 333 and 373 K, respectively.2542 SELF-REACTIONS OF ISOPROPYLPEROXY RADICALS Reaction (5) is one of the most important in the reaction scheme, the value assigned to k, having a profound effect on the simulated product distributions. Its importance is particularly apparent when the effect of varying oxygen pressure is considered. We have used the Arrhenius parameters from the corresponding reaction between methoxy radical and oxygen [reaction (16)] which were obtained by Batt and Robinson;20 these explain successfully the results not only on varying pressures of oxygen isothermally, but also on altering the temperature.The only reported value for the corresponding reaction to (8) comes from a liquid-phase study2' of the radical-induced decomposition of several tertiary hydro- peroxides; a value for the rate constant for abstraction from hydroperoxides of the hydroperoxidic hydrogen by alkoxy radicals of ca. 4 x lo6 dm3 mol-l s-l at 303 K was suggested. From the present results it can readily be seen that attack on isopropyl hydroperoxide must be taking place and, moreover, that this attack is reduced by increasing the oxygen pressure, strongly suggesting that the attacking species is the isopropoxy radical.The absolute value of k, must be high enough to compete favourably with termination reactions (3a) and (7). Reaction (7) is the major alternative fate for isopropoxy radicals to reactions (5) and (8), and it is therefore the value assigned to reaction (7) that decides what absolute values of k, and k , will produce a good fit of simulated to experimental results. This, coupled with the ratio k,/k, obtained, leads to values of k, (and by analogy k,,, k,, and kZO) of 5.0 x lo7, 6.2 x lo7 and 7.7 x lo7 dm3 mol-l s-l at 302, 333 and 373 K. Following our estimates, Kirsch and Parkes have suggested that the rate constant for the reaction between alkoxy radicals and t-butyl hydroperoxide is ca.6 x lo7 dm3 mol-l s-l between 298 and 373 K.2 This indicates that the reaction is faster than the reaction in which an alkoxy radical abstracts a tertiary hydrogen from 2-methylpr0pane.~~ This would agree with the proposition that the activation energy for the following general reaction XH+*Y +X---H---Y +Xe+YH can be represented by the difference between the energies of the X-H bond and the summation of the energies of the X- - -H and Y- - -H The total bond energy of the three-body transition state is thus dependent upon the X---H and Y---H distances and the values for the bond dissociation energies of X-H and Y-H. If we now consider the following reactions RH+RO* +Re +ROH R0,H + ROO -+ RO; + ROH it can be seen that given similar transition states the magnitude of the activation energies will be determined essentially by the bond dissociation energies ( D e ) of ROO-H and R-H.The bond dissociation energy of ROO-H of ca. 366 _+ 8 kJ mol-l can be calculated based on the assumption that the bond dissociation energy of ROO-H is the same as the bond dissociation energy of H00-H.25 A value for the latter is obtained from a value for the standard heats of formation of the hydroperoxyl radicalZs and the standard heats of formation of hydrogen peroxide and hydrogen atoms. D e (ROO-H) thus is much lower than the 0-H bond dissociation energy in other molecules. For example, in the case of t-butyl alcoholZ7 the value is 440 kJ mol-l.It is also lower than that accepted for the C-H bond dissociation energy in 2-methylpr0pane,~~ 380.4 kJ mold', which explains the results from independent experiments we carried out. 2-Methylpropane and t-butyl hydroperoxide were added to 2,2'-azopropane + oxygen mixtures and photolysed. The experiments show con- clusively that the hydroperoxide is consumed much more rapidly than the hydro-L. T. COWLEY, D. J. WADDINGTON AND A. WOOLLEY 2543 carbon,27 confirming that the activation energy for hydrogen abstraction by isopropoxy radicals from 2-methylpropane is significantly larger than that for abstraction of hydrogen, R02-H, by isopropoxy radicals from alkyl hydroperoxides. The corre- lation between the activation energy for hydrogen abstraction and the strength of the original H-X bond, when the attacking species are the same, seems very reasonable.The success of the calculations of Z a v i t s a ~ ~ * , ~ ~ adds considerable weight to such assumptions. To confirm that the alternative abstraction from a C-H bond in the hydroperoxide was unlikely, as suggested by considering the bond dissociation energies, the reaction was added to the simulations. The reaction yields hydroxy radicals, and these in turn react with reactant and products. Indeed it was not possible to simulate the reaction successfully. The mechanism predicts the formation of small amounts of methyl hydroperoxide. However, even under the most favourable conditions for its formation, it would account for less than 1% of total products. The failure to detect this ubiquitous compound is not a serious indictment of the proposed mechanism. At high oxygen concentrations, even at 373 K, the reaction is much simplified, as the formation of acetaldehyde, formaldehyde and methyl alcohol are supressed. Thus the rate of reaction ( 5 ) is much faster than that of reactions (4), (9, (9) and (1 l), and in the early stages of reaction2 Values for this ratio were found by determining the rates of formation of acetone and isopropyl alcohol at low extents of photolysis in the presence of excess oxygen (table 1). Using these results, the rate constants and Arrhenius parameters for reactions (3a) and (3b) are obtained (tables 1 and 2).TABLE VALUES OF RATE CONSTANTS FOR THE REACTIONS 2(CH3),CH0, -, products (ow 2(CH3),CH0, -+ (CH3),C0 + (CH,),CHOH + 0, (3 a) 2(CH3),CHO; --* 2(CH3),CH0.+ 0, (3 b) (k3a+2k3d k36 kobsl k3aI k3d lo5 dm3 mol-l s-l lo6 dm3 mol-l s-l lo5 dm3 mol-l s-l TIK k3a k3a 302" 3.78f0.10 1.39k0.04 8.1 & 0.5 2.15 f 0.10 2.99 f 0.20 333 4.65 f 0.10 1.83 f 0.04 16.2 Ifr 1.2 3.5 & 0.7 6.4 f 1.3 373 6.60 k0.20 2.80 f 0.08 35.6 f 2.0 5.4+ 1.1 15.0f 3.0 a Obtained earlier.'? The assumptions made in finding k3, and k3b appear to be justified. kobs does not vary with oxygen pressure,l as predicted by the model, confirming that (k3,+2k3,) = kobs under the conditions of the experiments. At low pressures of oxygen, the rates of reactions (7) and (8) are significant. However, the rates of reactions ( 5 ) and (6) are reduced, the overall effect being that the observed rate constant of self-reactions, kobs, remains constant as the oxygen pressure is varied although the rates of formation of individual products change.No other data for the separate self-reactions of alkylperoxy radicals have been2544 SELF-REACTIONS OF ISOPROPYLPEROXY RADICALS TABLE 2.-vALUES OF ARRHENIUS PARAMETERS FOR THE REACTIONS 2(CH,),CHO, + products (obs) 2(CH&CHO, + (CH,),CO + (CH,),CHOH + 0, (3 4 2(CH3),CH0,* + 2(CH3),CHO* + 0, (3 b) ~~ reaction A/dm3 mol-l s-l E/kJ mol-1 obsa 1.43+0.10x lo9 18.7k0.5 3a 2.44f0.31 x lo7 12.0+ 1.0 3b 1 . 3 8 k 0 . 2 6 ~ lo8 21.3+ 1.5 a Obtained earlier.' reported. The Arrhenius parameters for the overall self-reaction in the liquid phase of but-2-yl-, cyclopentyl-, cyclohexyl-, phenylethyl- and hexadecyl-peroxy radicals have been determined,29 Aobs being in the range (1-10) x lo9 dm3 mol-l s-l and Eobs in the range 7-13 kJ mol-l.However, in a recent study of the self-reactions of isopropylperoxy radicals in the liquid phase, Furminsky et al. report values for rate constants between 175 and 229 K which are very close to the values predicted by the Arrhenius parameters given in this Before the Arrhenius parameters of the separate reactions are discussed in detail, the values for other systems are needed. In a subsequent paper, the reactions of ethylperoxy radicals are described. The lines drawn in fig. 2 and 3 are the results produced by computer simulation using the chemical model just described. The success of the simulation technique is illustrated in particular when the temperature is varied and when the initial oxygen pressure is altered.2 To show how quantitative conclusions concerning rate constants can be made using the simulation technique it is useful to consider the effect of changing the values of k, and k,.As discussed earlier, it is the ratio of these two rate constants which govern the change in the relative rates of formation of isopropyl alcohol and isopropyl hydroperoxide as the oxygen pressure is changed. If the rate constants are changed to say k, = 6 x 102 dm3 mol-1 s-l and k, = 1 x lo5 dm3 mol-l s-l at 302 K, values which give the required ratio of rate constants, the simulated results bear no resemblance to the experimental ones. However, if the value of k, is set at 1.84 x lo6 dm3 mol-1 s-l, thus returning to the original value of the constant (k, + k,)/k, and, in addition, if k,, k, and k,, are each given a value of 1 x lo6 dm3 mol-l s-l, then the simulation once more fits the experimental results well. The implication of this result is clear; if the ratio of k,/k, is kept at 166 and the experimentally determined rate constants, k,, k,, and k,, are untouched, then it is always possible to produce simulated results which match the experimental results well.However, this is only true if we abandon the chemical intuition which makes certain values for unmeasured rate constants acceptable. The arguments put forward when discussing the individual rate constants show that the fit between simulated and experimental results is obtained in this case only by ignoring large sections of chemical knowledge.This is clearly unacceptable when perfectly good simulations can be obtained if this knowledge and its subsequent implications are taken into account. It is with some justification, therefore, that quantitative limits are put on the absolute values of k, and k , using the evidence of the present work. If the chemical model is a valid one, the computer simulation should be capable of predicting the behaviour of radicals in the system. Fig. 4 illustrates that, indeed,L. T. COWLEY, D. J. WADDINGTON AND A. WOOLLEY 2545 0.3 c 0 0.2 5 -2 5 0.1 2.0 4.0 6-0 8.0 t l s FIG. 4.-Photo-oxidation of 2,2’-azopropane at 302 K. Changes in absorption due to isopropylperoxy radicals. 1.0 2.0 3.0 4.0 5.0 t Is FIG. 5.-Photo-oxidation of 2,2’-azopropane at 302 K.Second-order plot for the simulated decay of isopropylperoxy radicals. the chemical model described here is consistent with the kinetics of the reactions of isopropylperoxy radicals. The solid line in the figure is the calculated change in absorption due to isopropylperoxy radicals at 302 K. The calculation used the values of the rate constants described above. The value of k , in the computer simulation was changed to zero at a time corresponding to the experimental time at which the photolysis lamps were switched off. To convert the concentration profile obtained from the simulation to one of absorbance, the experimental path length and value for the absorption cro~s-section~~ of the isopropylperoxy radical were used (fig. 5), illustrating that if the chemical model proposed here is a true reflection of the experimental system then the isopropylperoxy radicals would display pure second-order kinetics.Simulations were also carried out which showed the effect of oxygen pressure on the rate at which isopropylperoxy radicals were formed and consumed. The2546 SELF-REACTIONS OF ISOPROPY LPEROXY RADICALS computed concentration of isopropylperoxy radicals, at any photolysis time, was changed on altering the oxygen pressure from 10 to 500 Torr by < 0.1 %. The effect is therefore negligible and the model is consistent with the observation that kobs does not vary experimentally with initial oxygen pressure. A. W. thanks the S.R.C. for a CASE studentship. We thank Dr C . Anastasi, Dr L. J. Kirsch and Dr D.A. Parkes (Shell Research Ltd) for invaluable discussions and Dr A. Prothero (also of Shell Research Ltd) for considerable assistance with the computations during the study. L. J. Kirsch, D. A. Parkes, D. J. Waddington and A. Woolley, J . Chem. SOC., Faraday Trans. 1,1978, 74, 2293. L. J. Kirsch, D. A. Parkes, D. J. Waddington and A. Woolley, J . Chem. Soc., Faraday Trans. 1,1979, 75, 2678. J. C. Stowell, J. Org. Chem., 1967, 32, 2360. J. I. Abram, G. S. Milne, B. S. Soloman and C. Steel, J. Am. Chem. SOC., 1969,91, 1220. G. 0. Pritchard and F. M. Servedio, Znt. J. Chem. Kinet., 1975, 7, 99. @ S. Chervinsky and I. Oref, J. Phys. Chem., 1977, 81, 1967. ' L. D. Fogel and C. Steel, J. Am. Chem. SOC., 1976, 98, 4859. @ P. Arrowsmith and L. J. Kirsch, J . Chem. SOC., Faraday Trans. 1, 1978, 74, 3016. lo A. Prothero, personal communication. l1 M. J. Yee Quee and J. C. J. Tynne, J. Phys. Chem., 1968, 72, 2824. l2 J. H. Ferguson and L. Phillips, J. Chem. SOC., 1965, 4416. l3 A. C. Baldwin, J. R. Barker, D. M. Golden and D. G. Hendry, J. Phys. Chem., 1977, 81, 2483. l4 D. L. Cox, R. A. Livermore and L. Phillips, J. Chem. SOC., 1966, 245. l5 L. Batt and R. T. Milne, Znt. J. Chem. Kinet., 1977, 9, 141; L. Batt, Znt. J. Chem. Kinet., 1979, 11, R. Klein, M. D. Scheer and R. Kelley, J. Phys. Chem., 1964, 68, 598. 977. N. Y. Al-Akeel and D. J. Waddington, J. Chem. SOC., Perkin Trans. 2, to be published. l7 D. A. Parkes, Int. J . Chem. Kinet., 1977, 9, 451. l8 D. A. Parkes, 15th Znt. Symp. Combustion (The Combustion Institute, Pittsburg, 1975), p. 795. C. Anastasi, D. J. Waddington and A. Woolley, J . Chem. SOC., Faraday Trans. 1, submitted for publication. 2o L. Batt and G. N. Robinson, Int. J. Chem. Kinet., 1979, 11, 1045. 21 J. A. Howard and K. U. Ingold, Can. J. Chem., 1969, 47, 3797. 22 L. Kirsch and D. A. Parkes, J . Chem. SOC., Faraday Trans. 1, 1981,77, 293. 23 M. A. Sway and D. J. Waddington, to be published. 24 A. A. Zavitsas, J. Am. Chem. SOC., 1972, 94, 2779; A. A. Zavitsas and J. A. Pinto, J. Am. Chem. 25 S. W. Benson, J. Am. Chem. SOC., 1965, 87, 972. 2@ C. J. Howard, J. Am. Chem. SOC., 1980, 102, 6937. 27 A. Woolley, D. Phil. Thesis (University of York, 1980). 28 N. Y. Al-Akeel, K. Selby and D. J. Waddington, J. Chem. SOC., Perkin Trans. 2, 1981, 1036. 2@ J. E. Bennett, unpublished results. 30 E. Furminsky, J. A. Howard and J. Selwyn, Can. J. Chem., 1980, 58, 677. SOC., 1972, 94, 7390; A. A. Zavitsas and A. A. Melikian, J. Am. Chem. SOC., 1975, 97, 2757. (PAPER 1 / 183 1)

ISSN:0300-9599    

DOI:10.1039/F19827802535

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1982, 78, 2547-2552 Electron Spin Resonance Studies of y-irradiated Lithium Alkyls The Methylene Anion7 BY MARTYN C. R. SYMONS* AND MILDRED M. MAGUIRE~ Department of Chemistry, The University, Leicester LE1 7RH Received 30th November, 198 1 Exposure of methyl lithium to “OCo y-rays at 77 K gave a radical having a 1 : 2: 1 triplet e.s.r. spectrum characteristic of *CH; anions. The perdeuterated compound gave a singlet due to .CD; anions. Similarly, ethyl lithium gave a quintet due to (CH$H)- anions and isobutyl lithium gave a poorly reselved multiplet assigned to EtcMe- anions. No evidence for the tautomeric change (CH,cH)- + (CH,=CH,)- could be obtained. The fall in the 8-proton hyperfine coupling on going from CH,CH, (26 G) to CH,CH- (21 G) is explained in terms of a decrease in hyperconjugation.Although the methylene molecule has been extensively studied by e.s.r. spectroscopy,l* we know of no authentic e.s.r. studies of the methylene anion, *CH,, or of its alkyl derivatives, *CR;. These anions have recently been invoked as significant intermediates in reactions of vic dihalide~.~ In our search for CR; anions, we have studied the effect of ionizing radiation on 1,l -dihalides but no species with the expected properties was ~btained.~ We considered the possibility of hal; formation: R,C hal, + e- + R,C: + hal; (1) but concluded that outer features attributable to hyperfine coupling to two equivalent halogen nuclei were not due to hal; ions but were probably due to (R,C hal,)+ ions.4 The major electron capture process was always (2) R,C hal, + e- -+ R,c ha1 + hal- the R,c ha1 radicals, whose e.s.r.spectra are quite character is ti^,^? showing no tendency to undergo further reduction. Irradiation of di-iodomethane adsorbed on silica gel gave a species with the expected 1 :2: 1 triplet e.s.r. spectrum, but this was firmly assigned to H,c-OSie radicals rather than *CH;.’ We therefore concluded that it would be better to start with (R,CH)- anions as precursors, and we selected lithium alkyls as a good source. Our results suggest that we have been successful. Radiolytic studies of metal alkyls have been quite extensive, but we know of no previous work on the lithium salts. In our own work on the radiolysis of dialkyl mercury and alkyl mercury halides,** electron addition gave the parent anions or, for the halides, dissociative electron capture, and electron loss, followed by proton loss, gave a- or B-mercury derivatives. The radical H,e-HgMe t Unstable Intermediates, Part 201.$ Present address : Department of Chemistry, Waynesburg College, Waynesburg, Pennsylvania 15370, U.S.A. 25472548 E.S.R. STUDIES OF LITHIUM ALKYLS had an e.s.r. spectrum with a 1 : 2: 1 triplet similar to that expected for *CH; anions, but well defined satellite lines from species containing lQQHg or ,OIHg clearly established the presence of the covalent bond to the HgMe group. Finally, it is interesting to note that the radical cation, *CHz was postulated many years ago as a product of the radiolysis of methanol and related compounds.lo Our studies on photolysed systems containing hydrogen peroxide, however, established that the species detected were, in fact, H,cOH radicals,ll as has been fully sub- stantiated.12 EXPERIMENTAL Methyl lithium was obtained commercially (Pfaltz and Bauer) as a 1 mol dm-3 solution in diethyl ether.Ethyl lithium (Alfa Ventron) was obtained as a 1 mol dm-3 solution in benzene. s-Butyl lithium (Alfa Ventron) was obtained as a 1 mol dm-3 solution in cyclohexane. Lithium hydride and lithium metal were obtained commercially (B.D.H. Chemicals). [2H3] Iodomethane, 99 atom % D, Gold Label, was also obtained commercially (Aldrich Chemicals Ltd).13 Solvents were removed from the alkyl lithium samples under vacuum, care being taken to exclude oxygen, and e.s.r. analysis and y-radiolyses were conducted in Suprasil quartz tubes sealed in uacuo.All samples were handled in a dry, oxygen free, nitrogen atmosphere. To remove unwanted glass resonances from the evacuated sample tubes, after irradiation the upper portions of the sealed tubes were flamed and the tube was inverted after cooling so that the samples were contained in the flamed portion of the tubes. Samples were irradiated in a Vickrad 6oCo y-ray source at a dose rate of ca. 1 Mrad h-l at 77 K for up to 2 h. The e.s.r. spectra were obtained with a Varian E-109 spectrometer equipped with 100 kHz field modulation, a Hewlett-Packard 5255A frequency converter and a Varian E-935 EPR bats acquisition system. All e.s.r. measurements were recorded at 77 K. Spin resonance g-values were obtained using a field marker of y-irradiated Suprasil whose g-value had previously been ascertained with respect to DPPH.Q-band measurements at 77 K were made with a balanced bridge reflection spectrometer constructed in these laboratories. RESULTS AND DISCUSSION E.S.R. SPECTRA AND IDENTIFICATION The spectra are shown in fig. 1-3, the data being summarised in table 1, together with data for *NH, and CH,NH radicals. Comparison of fig. 1 and 2 shows conclusively that the 1 : 2: 1 triplet obtained from LiMe is due to hyperfine coupling to two equivalent protons. There is no resolved coupling to sLi or 'Li so that this interaction must be < 3 G. Although the three features are broad, there is no clear evidence for g-anisotropy. Coupling to ,H is not resolved for the species formed in LiCD,.The isotropic coupling for 2H obtained from the lH coupling of ca. 20 G must be ca. 3 G, which is comparable to the expected coupling to lithium nuclei. We suggest that the unusual shape of the singlet shown in fig. 2 arises because certain features of the expected multiplet are better defined than others. This species is presumably H,C-/Li+. The anion CH; is isoelectronic with the -NH2 radical which has an isotropic proton coupling of ca. (-) 24 G, so our value of ca. (-) 20 G is quite reasonable. Lithium alkyls show a strong tendency to form cluster~,~*9 l5 but the metal-carbon bond is certainly strongly ionic. Hence, to a first approximation, we favour the ion-pair formulation for H,t]/Li+. Thiscan be compared, for example, with the ion-pair 0,C *-/Li+ formed in y-irradiated lithium formate.16 The coupling to 6v7Li nuclei in this species is 3.3k0.2 G, which is certainly greater than any coupling that may be present in our system. A small cation hyperfine coupling accords with the structure shown in (I).AsM. C. R. SYMONS A N D M. M. MAGUIRE 2549 c 3200G ( 8.969GHz 1 FIG. I.-First derivative X-band e.s.r. spectrum for CH,Li after exposure to 6oCo y-rays at 77 K, showing features assigned to H,t-/Li+ ion-pairs. suggested previously,ls9 l7 coupling for this structure will be via spin-polarisation of the lone-pair electrons and should, therefore, be negative. However, libration of the cation relative to the anion will move it out of the plane towards the orbital of the unpaired electron, which will introduce a positive coupling due to direct delocalisation.Z I I The small resultant coupling stems from partial cancellation of these two contributions. For CO; radicals in a range of salts, alkali-metal cations modify the g- and A-tensor components only to a small extent,16 and hence the present data assigned to *CH; radicals are expected to be quite close to those for the unperturbed anions. THE g-VALUE Expectation for *CH; radicals would be for g, = 2.0023, with g, 2.003 and g, greater than this. This is because g, lies along the axis of the half-filled orbital, g, has a contribution from coupling with the a-electron levels and hence should be almost equal to that for *CH, (ca 2.003), but g, includes coupling with the weakly bonding ‘lone pair’, which should contribute a greater positive shift.In fact, there is no clear g-anisotropy, the limiting range imposed by the central component in fig. 1 being ca. 0.005 so that if g, = 2.0023, g, < 2.007. Unfortunately, even at Q-band frequencies2550 E.S.R. STUDIES OF LITHIUM ALKYLS I 32306 ( 9.053GH~ 1 FIG. 2.-As for fig. 1, showing features for D$-/Li+ formed from CD,Li. I 32506 (9-109GH~) I I - 2 1 1 ) - 2 w I - 1 FIG. 3.-As for fig. 1, showing features for CH$H-/Li+ formed from CH,CH,Li. the expected anisotropy was not well defined, the results setting an upper limit of CQ. 2.005 for g,. A similar situation is found for trapped *NH, radicals.le No clear g-anisotropy is detected, the mean value from the non-rotating radical being close to gav of ca.2.005 obtained from rotating radicals. However, the g-tensor components obtained from laser magnetic resonance studies of *NH, in the gas phaselo (gz = 2.0023, g, = 2.0040, gy = 2.0088 on the previous assignment) fit in well with expectation. Again, anisotropy is not well defined for the lH coupling in the solid-state spectra. This arises in part because the turning points for the combined splitting will lie alongM. C. R. SYMONS A N D M. M. MAGUIRE 2551 TABLE ELECTRON SPIN RESONANCE DATA hyperfine components/Ga H' (C) H$-/Li+ 24.5 19 ca. 20 - 2.005 - - 2.005 CH$H-/Li+ (aH) 26 10 ca. 20 - 2.006 ca. 21 *NH,b - - - 16.4, 27.1, 28.6 2.0088 (g,) 2.0040 (gz) MeNHC - - 33.0 26 2.002 (min) 2.005 (max) D$'-/Li+ - - - - - (BH) - 2.0023 (g,) a G = T; * ref. (19); ref. (21).the symmetry axes rather than along the C-H bond directions so that the full anisotropy is not obtained. As can be seen from fig. 3, this is not the case for MecH- radicals since the form of the spectrum is now controlled by the anisotropy of the single a-proton. The isotropic coupling of ca. 20 G is ca. 10% less than that for *CH,. A small reduction is expected by comparison with results for *NHZ (ca. 25.8 G) and *NH, (ca. 24G). THE RADICAL (CH,CH)- 1 lines showing the expected anisotropy. The central component is stronger than expected and probably includes a singlet from some unidentified species. The mean hyperfine coupling is ca. 21 G. We suggest that the radical is (CH,CH)- in the ion-pair CH,CH-/Li+, rather than the tautomer (CH,-T-CH,)-. The ethylene radical anion has never been studied by e.s.r.spectroscopy, so far as we know, but simple expectation would suggest that the proton coupling should be ca. 1 1 G, since each carbon can only have a spin density of 0.5. If the radical is not planar at carbon the magnitude of the lH coupling will be less than ca. 1 1 G. This is supported by reference to the isoelectronic radical N,H; that we studied some time ago.2o The isotropic proton coupling was ca. (-) 1 1 G. Since the present coupling is ca. 21 G the species is surely CH,cH-/Li+. If there were any tendency for a tautomeric change, it would be opposed by the Li+ ion interacting with the lone pair of electrons. The anion CH,CH- is expected to have an e.s.r. spectrum comprising a nearly isotropic 1 : 3: 3: 1 quartet from the CH, group (which is expected to be rotating by analogy with CH,CH, radicals trapped at 77 K), each component being split into an anisotropic doublet from the single a-proton.From the anisotropy we find that the maximum coupling to a-H is ca. 26 G and the minimum is ca. 10 G measured from the - 2 and - 1 components. If we assume that Aiso z 20 G as for CH, radicals, then the three components are 26, 24 and 10 G. These values can be compared with those for *NH, radicals in the gas phase (28.6, 27.1 and 16.4). The anisotropic contributions are quite similar, being slightly larger for the c-H unit, as expected. The value of ca. 21 G for the three P-protons should be compared with that of ca. 26G for CH,tH, and ca. 33G for CH,NH.,l The marked reduction can be The spectrum (fig.3) comprises a set of five hyperfine components, the It 2 and2552 E.S.R. STUDIES OF LITHIUM ALKYLS understood in terms of the concept of hyperconjugation, which has been shown to be a function of the effective electronegativity of the atom on which the unpaired electron is formally localised.22123 On going from CH,cH, to CH,cH- there is a marked fall in effective electronegativity and hence a fall in the #?-proton coupling. This concept explains the steady fall from CH,NH to CH,cH, to CH,CH . ISOBUTYL LITHIUM Our results for this compound were disappointing. Weak broad lines were obtained in the regions expected for the species (CH,CH,cCH,)-, but they were partially concealed by an intense central component from unidentified radicals.All we can conclude is that a p-proton coupling of 20-25 G is required to fit the outer components. Probably the coupling constants to the CH, protons differs from that to the CH, protons, implying some degree of restricted rotation for the ethyl group. M.M.M. thanks the Leverhulme Trust of the United Kingdom for the award of a Leverhulme Visiting Fellowship and also the American Association of University Women for the award of a Curie Fellowship. R. A. Bernheim, H. W. Bernard, P. S. Wang, L. S. Wood and P. S. Skell, J. Chem. Phys., 1971, 54, 3223. E. Wasserman, W. A. Yager and V. J. Kuck, Chem. Phys. Lett., 1970, 36, 409. G. D. Sargent, C. M. Tatum and R. P. Scott, J. Am. Chem. Soc., 1974, 96, 1602. S. P. Mishra and M. C. R. Symons, J. Chem. SOC., Perkin Trans.2, 1975, 1492. S. P. Mishra, G. W. Neilson and M. C. R. Symons, J. Chem. SOC., Faraday Trans. 2, 1973,69, 1425; S. P. Mishra, G. W. Neilson and M. C. R. Symons, J. Chem. SOC., Faraday Trans. 2, 1974,70, 1165. K. Shimokoshi, Bull. Chem. SOC. Jpn, 1974, 47, 11. ' B. W. Fullam and M. C. R. Symons, J. Chem. SOC., Dalton Trans., 1974, 1086. * M. C. R. Symons and M. M. Aly, J. Organomet. Chem., 1979, 166, 101. W. Gordy and C. G. McCormick, J. Am. Chem. SOC., 1956,78,3243; J. Luck and W. Gordy, J. Am. Chem. SOC., 1959, 81, 269. J. F. Gibson, D. J. E. Ingram, M. C. R. Symons and M. G. Townsend, Trans. Faraday SOC., 1957,53, 914. lo J. F. Gibson, M. C. R. Symons and M. G. Townsend, J. Chem. SOC., 1959, 269. l2 W. T. Dixon and R. 0. C. Norman, J. Chem. SOC., 1963, 31 19. l3 H. A. Gilman, E. A. Zoellner and W. M. Selby, J. Am. Chem. SOC., 1932, 54, 1957; H. A. Gilman, l4 G. Graham, S. Richtmeier and D. Dixon, J. Am. Chem. SOC., 1980, 102, 5759. l5 R. Zergec, W. Rhine and G. Stucky, J. Am. Chem. SOC., 1974, 96, 6048. J. Am. Chem. SOC., 1933, 55, 1253. J. H. Sharp and M. C. R. Symons, J. Chem. Sac. A , 1970, 3075. M: C. R. Symons, Nature (London), 1969, 224, 685. I. S. Ginns and M. C. R. Symons, J. Chem. SOC., Faraday Trans. I , 1972, 68, 631. l9 P. B. Davies, D. K. Russell, B. A. Thrush and H. E. Radford, Chem. Phys. Letr., 1976,42, 35. 2o J. A. Brivati, J. M. Gross, M. C. R. Symons and D. J. A. Tinling, J. Chem. SOC., 1965, 6504. M. C. R. Symons, J. Chem. SOC., Perkin Trans. 2, 1973, 797. 22 J. A. Brivati, R. Hulme and M. C. R. Symons, Proc. Chem. Soc., 1961, 384. 23 R. Hulme and M. C. R. Symons, J. Chem. SOC., 1965, 1120. (PAPER 1 / 1859)

ISSN:0300-9599    

DOI:10.1039/F19827802547

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Furuduy Trans. 1, 1982, 78, 2553-2561 Thermal Decomposition of Ammonium Chromate BY SUNDARA R A J A M ~ AND ANDREW K. GALWEY* Department of Chemistry, The Queen’s University of Belfast, Belfast BT9 SAG Received 30th November, 198 1 Kinetic evidence has shown that the decomposition of ammonium chromate [+ (NH,),Cr,O, + H20 + 2NHJ obeys the contracting cylinder equation. Evidence that a reaction interface advances progressively inwards from the original crystal surfaces was obtained from microscopic examination of cross-section surfaces revealed by fracture of the lath-shaped crystals of reactant after partial decomposition. Consistent and complementary results were obtained from quantitative studies using three different salt preparations. It is concluded that two distinct chemical steps contributed to reaction.These were reversible proton transfer, which was opposed by higher pressures of accumulated product gas, followed by water elimination, which was not reversible under reaction conditions. The decomposition is very slightly endothermic but is irreversible, an unusual combination. Although there have been numerous kinetic. and mechanistic studies of the decomposition reactions of solids,l at present there are no generally accepted criteria for the classification of such rate processes into groups exhibiting similarities in patterns of chemical, kinetic or mechanistic behaviour. Garner, in his important survey, simply made a distinction between endothermic2 and exothermic3 reactions. Endothermic rate processes are usually reversible and typically involve the release of ligands and/or molecules loosely bonded within the crystal structure, as, for example, in the release of water from crystalline hydrates.In contrast, exothermic processes are frequently irreversible, since one or more of the crystal constituents undergoes primary bond reorganization, as in the breakdown of anions in salts of oxyacids. Similarities between reactions for groups of salts containing common or related constitutents have been described and discussed.’ To investigate the general applicability of Garner’s classification, we have studied the mechanism of the decomposition of ammonium chromate4 2(NH4),Cr04 + (NH,),Cr,O, + 2NH, + H 2 0 since this reaction exhibits characteristics of both of the classes distinguished. The chemical change is reversible and the volatile products are molecules that are frequently found as ligands released in endothermic reactions [eqn (l)].However, the cation-anion interaction results in a primary bond redistribution involving the anion, a feature often found in irreversible reactions, though here it is probably significant that the oxidation state of the chromium (CrV1) remained unchanged. Another reason for our study was that although the decomposition of ammonium dichromate has been the subject of several studies,*-’ much less interest has been shown in this closely related, indeed precursor, rate process. The reaction of (NH4),Cr04, however, contrasts with that of (NH,),Cr,O, in that ammonia is evolved in the t On leave from Q.E.M.Government College for Women, Madras 600002, India. 25532554 THERMAL DECOMPOSITION OF AMMONIUM CHROMATE lower-temperature reaction of the former salt whereas decomposition of the latter is accompanied by partial oxidation of the constituent nitrogen to N, and N,O. EXPERIMENTAL AMMONIUM CHROMATE B.D.H. Laboratory Reagent ammonium chromate was used throughout the work. The composition determined by combustion analysis (18.35% N and 5.20% H, also4 < 0.1% C) agreed well with theoretical requirements for (NH4),Cr04 (18.42% N and 5.30% H). Reactant samples were stored for the minimum time practicable in an atmosphere of moist ammonia to minimize the effects of deterioration due to ammonia loss. Three samples of (NH4),Cr04 were investigated. The purchased salt is referred to as the ‘crystallites’.This material was recrystallized from water containing excess ammonia to yield the ‘crystals ’. Some crystals were compressed (ca. 100 MN m-,) to form a single coherent ‘pellet’, individual fragments of which were used in rate studies. APPARATUS Kinetic measurements were made using a constant-volume system in which the pressure of ammonia evolved by salt decomposition was determined at known times. The McLeod pressure gauge was separated by a cold trap (ca. 210 K, to condense water) from the reactant, which was held at constant temperature ( & 0.5 K). Initial sample outgassing (ca. Torr for 30 min) was minimized because salt breakdown occurred at ambient temperature. Scanning electron micrographs were obtained using a Cambridge Stereoscan S 180 : before examination samples were coated with gold/palladium.RESULTS AND DISCUSSION STOICHIOMETRY The mean weight of the residual solid, on completion of kinetic measurements, was 83 %, which agrees well with theoretical expectation (82.9%) assuming that reaction yields (NH4),Cr20, only. Most of the 25 values measured were between 82 and 84%. The quantities of gas evolved on completion of reaction, however, tended to be less than the theoretical quantity (often x O . 9 ) : this is attributed to losses during preliminary evacuation. Accordingly, the period of outgassing was reduced to a minimum and the significance of this loss on the kinetic analysis is considered below. Product gases were fully condensed at 78 K, showing that decomposition involved no appreciable (< 0.2%) oxidation of ammonia to nitrogen.s The stoichiometry of the reaction is, therefore, satisfactorily expressed by eqn (1).PHYSICAL MEASUREMENTS Differential scanning calorimetric measurements, heating at 20 K min-l, showed that the reaction was slightly endothermic: 13 J g-l or 2.0 kJ mo1-l. The surface area of salt previously decomposed in vacuum (2 h at 373 K), measured from adsorption of nitrogen at 78 K and application of the B.E.T. equation, was 2-3 m2 g-l. This area corresponds to an assemblage of uniform equal sized cubes, each of ca. 1 pm edge. KINETIC STUDIES CRYSTALS This preparation was studied here in the greatest detail since the large, lath-like crystals (ca. 1 x 1 x 10 mm) were well defined and particularly suitable for both kinetic and electron microscopic investigations.Some 30 isothermal kinetic measurementsS. RAJAM AND A. K. GALWEY 2555 tlmin FIG. 1 .-Typical fractional decomposition (a) against time plots for the decomposition of ammonium chromate crystals at four temperatures. 0.8 ’ - O l I 0.6 -5 I “ 1 I n r 0 20 30 i 0 50 60 70 t/min FIG. 2.4bedience of results given in fig. 1 to the contracting cylinder equation’ [ 1 -( 1 -a)& = k r]. The initial acceleratory process appearing in the experiment at 381 K is attributed to decomposition before attainment of thermal equilibrium and the later decrease in rate of the reaction at 356 K is ascribed to inhibition by ammonia gas (see text). were made between 355 and 382 K, with additional rate observations, concerned with the onset of reactions, from 334 to 355 K.Typical fractional decomposition (a) against time plots for the decomposition of crystals are shown in fig. 1. Reactions were deceleratory throughout and rate changes fit the contracting cylinder equation1 (shown for the same experiments in (2) fig. 2) 1 -(I -or)+ = kt.2556 THERMAL DECOMPOSITION OF AMMONIUM CHROMATE This obedience was regarded as acceptable only after a critical consideration of the accuracy of fit of data to this and to appropriate alternative rate expressions.*? Note, however, that during the latter stages of reactions at lower temperatures and with higher reactant weights the rate processes became more deceleratory than the requirements of eqn (2) (see reaction at 356 K in fig.2). Low temperature (334-348 K) experiments gave no evidence of any initial acceler- atory process : reactions were deceleratory from their inception, 5 min after introduction of the sample into the heated zone. The linear kinetic plots in fig. 2, and indeed all others of the series, intercepted the time axis at t = 5+ 1 min, which is, therefore, identified as the warm-up period. The indications of an initial acceleratory process in the higher temperature reactions (e.g. at 381 K in fig. 1) can be ascribed to perceptible onset of decomposition before the attainment of thermal equilibrium, since extrapolations of the linear portions of such plots (fig. 2) also intercepted the time axis at t = 5 min. A comparative analysis of all kinetic data showed that at low temperature a against time plots were more deceleratory than the requirements of eqn (2).Deviations from linearity of 1 - (1 - a)$ against time plots occurred at a characteristic pressure (p,) that was temperature dependent. Measurements of pD values were sensitive to error and fig. 3 shows a linear plot of log (pD/mm) against (T/K)-l. Confirmation that the 0.5. 0.0. n E E 1 brJ .3 - 0 . 5 26 27 28 29 30 104 K/T FIG. 3.-Decomposition of ammonium chromate crystals. Plot of mean pressures at which deviation from obedience to eqn (2) becomes appreciable: plot of log(p,/mm) against (T/K)--l. accumulation of gaseous products caused this relative reduction in rate was obtained from experiments in which accumulated gas was intermittently withdrawn. During a continued reaction at 353 K the product gases were evacuated briefly (3 min) at a = 0.20,0.46 and 0.69.Rates of subsequent accumulatory pressure rise (without any further 5 min induction periods) were always greater than those immediately prior to gas withdrawals and satisfied eqn (2). The same result was obtained from further similar experiments. There was a tendency for values of k to decrease with increasing reactant weight;S. RAJAM A N D A. K. GALWEY 2557 for example, in 7 experiments at 381 K the magnitude of k systematically decreased from 0.045 min-l, when < 10 mg crystal was used, to 0.035 min-l for 30-40 mg reactants. A similar spread of values was apparent throughout the temperature interval investigated in kinetic studies, 334-383 K. Such variations are consistent with the inevitable use of finer crystals in samples of smaller mass, since k = k,/a, where k , is the rate of interface advance and a is the cylinder radius.l Crushing reactant crystals also markedly increased decomposition rates, by as much as a factor of 10, attributable to an increase in reactant area and a decrease in the diameter of the reactant particles.From 33 values of k , measured between 334 and 383 K, the apparent activation energy, calculated by the root-mean-square regression analysis, was 97 f 5 kJ mol-l and the frequency factor (A/min-l) was log (A/min-l) = 12.05 f0.73. Exposure of partially decomposed salt to ca. 4 Torr 2NH,+H,O at 365 K under reaction conditions resulted in no appreciable gas uptake in 3 days. Thus, the observed relative decrease in decomposition rate at product pressures greater than p,, cannot be attributed to onset of the reverse reaction but may be due to effective opposition to the dissociation step due to the presence of products.At higher product pressures (> 20 Torr) and at ambient temperature, however, reacted crystals regained their original weights ( f 5 %) within 12 h, showing the reaction to be reversed [eqn (l)] under these changed conditions. The kinetic characteristics of decomposition of the reamminated salt were similar to the pattern of behaviour described above: a similar yield of ammonia was evolved by a rate process obeying eqn (2) between 0.03 < a < 0.95. However, reaction rates were greater, values of k being increased usually by factors of 7 to 10. For example, with one typical sample the almost equal rates of second and third deamminations were greater than the first by factors of 8.5 and 9.0, respectively. Crushing the reamminated material caused a further small increase ( x 1.5) in the rate of a subsequent decomposition.Some features of the decomposition of reamminated crystals are shown in fig. 4, an a against time plot for reaction at 356 K. After partial reaction (a = 0.22 at 28 min, point A) the salt was reamminated overnight in ammonia and saturated water vapour at ca. 280 K. After brief outgassing, the initial rate of subsequent decomposition was relatively rapid, identified as ammonia loss from the reamminated material, before decreasing to a rate which extrapolated as a continuation of the interrupted first reaction.(Point B, fig. 4, has been selected to emphasize this kinetic continuity of the first reaction as it proceeds towards completion.) MICROSCOPIC STUDIES The microscopic studies were largely directed towards characterization of the well defined textural changes apparent during deammination of the crystals. Similar observations for the other preparations were less easily interpreted, due to greater complexity of reactant textures, but all results for these supported the conclusions reached from examinations of the crystals. To permit meaningful comparisons, kinetic and scanning electron microscope examinations were concerned with similar specimens : lath-shaped crystals, 4- 10 mm in length and square (ca. 1.0 mm edge) cross-section with rounded corners.Planar surfaces of the unreacted crystals were usually smooth, though these sometimes contained adherent or embedded smaller crystals. After decomposition, numerous small polygonal, sometimes diamond-shaped, depressions appeared. Within each depression there was irregular crack development, infrequently containing a spherical particle : typical features are shown in plates 1 and 2. These pits and cracks are similar,2558 THERMAL DECOMPOSITION OF AMMONIUM CHROMATE - -.-- 0 10 i0 i 0 50 6-0 7-0 tlmin FIG. 4.-An ammonium chromate crystal was deamminated at 356 K to a = 0.22, point A, reamminated with water vapour and ammonia gas at ambient temperature before conditions of former decomposition were restored, point B. The initial loss of ammonia was rapid, attributed to losses from reamminated material, and subsequently the reduction in rate corresponds to continuation of the earlier reaction, identified as decomposition of unchanged reactant.(Point B was selected to emphasize the continuity of the reaction.) in some respects, to the appearance of nuclei developed during chrome alum dehydrationlo and here may indicate an initial and dense nucleation which results in rapid completion (a < 0.01) of surface reaction. Surfaces exposed on the transverse fracture of partially decomposed crystals showed that crack systems penetrated the outer layers of irregular particles, identified as product, and there was an inner zone of smoother cleavage, identified as reactant. This structure is apparent in the representative plates 3 (where a = 0.12 at 338 K) and 4 (where a = 0.17 at 348 K).No regular crack structure or other textural feature useful in elucidating the reaction mechanism could be discerned within the reactant-product contact zone at the highest magnifications available. The kinetic and microscopic observations are, therefore, complementary and entirely consistent. The reaction is initiated rapidly at all surfaces of the lath-shaped crystal and then advances inwards. The outer product layer is permeated by a system of fine cracks, permitting escape of the gases released. At higher maintained pressures of product (> pD) the water and ammonia retained within these channels may be sufficient to influence equilibria at the interface and so diminish the overall rate of product release.Nevertheless, when a cold trap ensures the irreversible removal of water, the reaction proceeds to completion. Reammination resulted in extensive retexturing of surfaces, to give, in some areas, the appearance of a mass of fibres (plate 5 ) while elsewhere crystallite development was more localized (plate 6). The present work did not examine this complex behaviour in detail but confirms that reammination results in reactant mobility leading to a large surface area increase and crack structure reorganization. These observations provide a satisfactory explanation of the relatively much more rapid rate of a second or third decomposition. Both reammination and crushing increase the decomposition rate ofJ. Chem. SOC., Faraday Trans.I , Vol. 78, part 8 Plates 1 and 2 Plate 1. Plate 2. PLATES 1 AND 2.-Typical textures of surfaces of ammonium chromate crystals after partial decomposition. The smooth surface of the reactant has become pitted and cracks have developed within such shallow pits. Crystals remained coherent on completion (a = 1 .OO) of decomposition. Distance between scale points at lower edge of each photograph is reported in the lower panel (also applies to plates 3-6). S. RAJAM AND A. K. GALWEY (Facing p . 2558)J . Chem. SOC., Faraday Trans. I , Vol. 78, part 8 Plates 3 and 4 Plate 3. Plate 4. PLAT= 3 AND 4.-Typical textures.of surfaces exposed on transverse fracture of lath-shaped ammonium chromate crystals after partial decomposition (plate 3, a = 0.12 at 338 K; plate 4, a = 0.17 at 348 K).Reaction has resulted in the development of an outer layer of crack-penetrated irregular material, identified as product, (NH,),Cr,O,, whereas the inner zone exhibited more regular, smoother cleavage and is identified as reactant. This is strong evidence of the operation of a contracting envelope mechanism. No detailed evidence concerning interface structure could be obtained. S. RAJAM AND A. K. GALWEYJ . Chem. SOC., Faraday Trans. 1, Vol. 78, part 8 Plates 5 and 6 Plate 5. Plate 6. PLATES 5 AND 6.-Typical surface textures of decomposed salt, reamminated at ambient temperature, giving evidence of extensive surface reorganization and acicular crystal growth. The extent of such modification varied between different areas of the surface.S . RAJAM AND A. K. GALWEYS. RAJAM A N D A. K. GALWEY 2559 crystals by an order of magnitude and both treatments cause a large increase in surface area of the reactant by the opening of internal surfaces and crack systems. CRYSTALLITES Decomposition characteristics of the crystallites resembled the behaviour of the crystals, though reaction was more rapid. Again eqn (2) was obeyed initially and values of k were greater, on average x 6, with a similar range of relative variation. Microscopic measurements gave the mean dimensions of the lath-shaped crystallites as 0.14 mm edge and 0.7 mm length. The mean effective edges of crystals and crystallites were therefore in the approximate ratio (1 .O/O. 14) NN 7 and this is (within experimental error for particle size measurement) in good agreement with the theoretical expectation that values of k in eqn (2) are inversely proportional to the particle radius (k = k,/a). From 14 values of k measured between 342 and 365 K., the calculated apparent activation energy was 90k9 kJ mol-l and log (Almin-l) = 11.94f 1.33.Both values agree with the results for crystals. While the measured values of pD for crystallites showed relatively greater scatter, mean magnitudes were comparable with the data for crystals (fig. 3). A sample of crystallites was crushed in a pestle and mortar and the influence of variation of sample weight on decomposition kinetics investigated at 345 & 1 K. Values of k systematically decreased from 0.024 min-l for 10 mg samples to 0.008 min-l for 40 mg samples (at the same temperature the mean k for crystallites was 0.008 min-l).This relative increase in rate ( x 3) of reaction of the 10 mg sample caused by crushing agrees well with theoretical expectation since optical measurements showed that average particle edge dimensions had been reduced from 0.14 to 0.05, corresponding to a change in a of x 1/2.8. The reduction in rate resulting from an increase in reactant weight is ascribed to the opposition to volatile product escape from between the particles of the more tightly packed product assemblages. Interruption of such reactions by a period of product evacuation resulted in a temporary increase in rate on resumption of accumulatory measurements. This is further confirmation that product gas accumulation within the reactant opposes progress of reversible decom- positions and aspects of the influence of such behaviour on kinetic characteristics have been discussed for nickel sq~arate.~ The value of k for a reaction at 360 K was 0.019 min-l and in three successive decompositions of the same salt, following reamminations, values of k were 0.076-0.08 1 min-l, also at 360 K.This relative increase in k ( x 4) is less than that found for crystals and pD values were greater for second and subsequent reactions. Both effects are readily attributed to the smaller particle size of this reactant, which was more reactive and less sensitive to product accumulation within the shorter pore systems. PELLET A single coherent pellet was prepared by compressing crystals ( 5 min at ca. 100 MN m-2) and 10 individual fragments used in rate studies.The data again obeyed eqn (2), and deviations at high a during reactions at low temperature occurred at pD values close to those given in fig. 3. At the upper end of the temperature range studied (358-390 K), values of k were close to those found for crystals, the edge dimensions of the reactant particles (1 .O-1.5 mm) being similar. At lower temperatures, however, pellet fragments reacted more rapidly than crystals for which the calculated apparent activation energy was 70 5 kJ mol-1 with log (A/min-l) = 8.5 0.5. The reason for this difference is not apparent but could be due to strain relief by aggregate2560 THERMAL DECOMPOSITION OF AMMONIUM CHROMATE disintegration during the longer period of reaction of the pellet at lower temperatures.This difference in behaviour emphasises the problems inherent in the measurement of meaningful values of activation energies for solid-state reactions and their use as a basis for formulating reaction mechanisms. CONCLUSIONS GEOMETRY OF REACTION INTERFACE DEVELOPMENT It is concluded from both kinetics and microscopic evidence that the decomposition of ammonium chromate can be represented'? * by the contracting cylinder equation, eqn (2). The reaction is initiated over all surfaces of the lath-shaped crystals and the reactant-product interface developed then advances inwards. Progressive changes of rate during decomposition of these asymmetric crystals can be regarded as reaction at the boundary of a contracting area, square or circle, both of which give eqn (2), neglecting the smaller contribution of end effects.Particle sections, from all our reactants, were approximately square, with rounded corners, and the term cylinder has been used above to emphasize the importance of the longer dimension in controlling kinetic behaviour. In eqn (2) k = k,/a, and the relative reaction rates ( k ) of crystals, crystallites and pellets (1 : 6: 1) were in agreement with this relation since values of a were in the ratio (1 : 1/7 : 1). Moreover, crushing also increased k values. This is further confirmation that eqn (2) provides a realistic representation of the reaction geometry, with further support from the microscopic evidence of the textures of cleaved partially reacted crystals (plates 3 and 4).The rapid initiation of reaction (fig. 1) is entirely consistent with obedience to eqn (2) and, indeed, the stoichiometric evidence is that some decomposition may have occurred during outgassing. Such losses would be expected to obscure the detection of any short initial acceleratory reaction period, which was sought but not found, though it was possibly indicated by microscopic evidence (plates 1 and 2) of a dense production of surface nuclei. Following superficial decomposition, subsequent reaction at a coherent interface proceeded at the boundaries of a cylinder of contracting diameter and eqn (2) was obeyed. REACTIONS AT THE REACTANT-PRODUCT INTERFACE The consideration of the changing geometry of the reaction interface, given above, does not provide evidence useful in the elucidation of the chemical steps occurring at that reaction interface.The formulation of a reaction mechanism must, therefore, be based on less direct evidence. The adherent layer of product (NH,),Cr,O, consists (plates 3 and 4) of irregular particles of perhaps 1pm edge. This is entirely consistent with the surface area measured for the residue and confirms the absence of any appreciable quantity of zeolite2 or other very finely divided product. The (NH,),Cr,O, layer produced is penetrated by narrow cracks (perhaps 0.1 pm width) affording escape routes for the evolved gases. No evidence of the development of a specialized interface structure was obtained from microscopic examinations. The observed behaviour can be satisfactorily accounted for by the following mechanism.On energetic grounds and by analogy with the decomposition of other ammonium salts,l it is reasonable to assume the first step to be reversible proton NH,+ + Cr0,2- $ NH, + HOCrO; transferS. RAJAM AND A. K . GALWEY followed by water elimination 2(HOCrO;) --+ H 2 0 + Cr20;- or, alternatively, there may be two proton transfer steps (HO),CrO, + CrOi- -+ H 2 0 + Cr20;-. 256 1 The kinetic observation that decomposition is opposed by higher pressures of ammonia is a consequence of the reversibility of proton transfer. The inhibitory effect increases with increasing amount of ammonia present and with reduction in reactant temperature, but such inhibition is rapidly removed on reaction vessel evacuation. The quantitative consideration of the kinetic consequences of participation of adsorbed species in surface reactions may be complicated.11 Accordingly we do not associate the calculated activation energy with a specific bond redistribution step. The decomposition is not reversible, as is indicated above, for the water elimination processes.Thus, although the reaction is slightly endothermic, it must be regarded as irreversible according to Garner’s classification.2T Following ammonia elimination, anion condensation requires extensive structural reorganization of the loosely packed, nearly tetrahedral CrOi- ions and NH; ions in (NH,),CrO, (monoclinic, space group C2/rn).l29l3 We have no evidence concerning the structure of the interface at which this change occurs. The reaction of residual (NH,),Cr,O, with saturated water vapour and ammonia at ambient temperature probably proceeds within an aqueous layer condensed within intercrystalline channels and covering all crystal surfaces.In these basic solutions, anion dimerization is reversed,14 dissolved chromate ions may migrate and recrystal- lization yields the elongated crystals on surfaces of reamminated salt (plates 5 and 6). Elongated crystallites were invariably formed on (NH,),CrO, recrystallization from aqueous solution. We conclude, therefore, that this reaction is not appropriately classified according to Garner’s scheme2- since it is (just) endothermic but not reversible under reaction conditions. The geometry of reaction interface development is identified as the contracting envelope from kinetic and microscopic evidence. Although a mechanism for the interface reactions, consistent with the available evidence, has been proposed, aspects of the chemical changes involved have not been fully characterized. M. E. Brown, D. Dollimore and A. K. Galwey, Comprehensive Chemical Kinetics (Elsevier, Amster- dam, 1980), vol. 22. W. E. Garner, Chemistry of the Solid State (Butterworths, London, 1955), chap. 8. W. E. Garner, Chemistry of the Solid State (Butterworths, London, 1955), chap. 9. I-H. Park, Bull. Chem. Soc. Jpn, 1972, 45, 2749, 2753. D. Taylor, J. Chem. Soc., 1955, 1033. J. Simpson, D. Taylor and D. M. W. Anderson, J. Chem. SOC., 1958, 2378. ’ B. Mahieu, D. J. Apers and P. C. Capron, J, Inorg. Nucl. Chem., 1971, 33, 2857. a M. E. Brown and A. K. Galwey, Thermochim. Acta, 1979, 29, 129. A. K. Galwey and M. E. Brown, J. Chem. Soc., Faraday Trans. 1, 1982, 78, 411. lo A. K. Galwey, R. Spinicci and G. G. T. Guarini, Proc. R. SOC. London, Ser. A, 1981, 378, 477. A. K. Galwey, Adv. Catal., 1977, 26, 247. l2 B. M. Gatehouse and P. Leverett, J. Chem. Soc. A , 1969, 1857. l 3 J. S. Stephens and D. W. J. Cruickshank, Acta Crystaliogr., Sect. B, 1970, 26, 437. l4 F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry (Interscience, New York, 2nd edn, 1966), p. 828. (PAPER 1/1860) 83

ISSN:0300-9599    

DOI:10.1039/F19827802553

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1982, 78, 2563-2571 Electrical and Electrochemical Properties of TiO, Films Grown by Organometallic Chemical Vapour Deposition BY YASUTAKA TAKAHASHI,* AKIFUMI OGISO, R ~ u z o TOMODA AND KOHZO SUGIYAMA Department of Synthetic Chemistry, Faculty of Engineering, Gifu University, Yanagido, Gifu 50 1 - 1 1, Japan AND HIDEKI MINOURA* AND MASAYASU TSUIKI Department of Industrial Chemistry, Faculty of Engineering, Gifu University, Yanagido, Gifu 50 1 - 1 1, Japan Received 7th December, 198 1 The electrical and photoelectrochemical properties of TiO, films, which were obtained by the vapour decomposition of ethyl titanate, are described. The electrical properties depend mainly on the deposition temperature. Films deposited at high temperatures have much lower electrical conductivity but remarkably higher photosensitivity than those deposited at low temperatures.The deposition temperatures are found to affect the electrochemical properties, such as anodic photocurrent-potential characteristics and spectra dependence of anodic photocurrents. Some doping effects of Al, Cr and Fe on the both properties are also examined. In a previous paper1 we described briefly the electrical and electrochemical properties of titanium dioxide (TiO,, rutile and anatase) films grown by the vapour decomposition of ethyl titanate. The as-grown films behave as a typical n-type semiconductor and have moderate electrical conductivity and large photoconductivity at room temperature without any additional treatment, such as reduction or heating following the deposition.TiO, is a stable photoanode, but can respond only to light which has a larger energy than its band gap (Eg = 3.0 eV), limiting its utility in an electrochemical photocell or in photocatalysis. Doping with various elements improves the response to ultraviolet or visible light.,-* We describe in this paper the details of an electrical and photoelectrochemical study of TiO, films obtained by our method, along with the effect of doping on the electrical and photoelectrochemical properties. EXPERIMENTAL The apparatus for the deposition of TiO, films and the starting complex, ethyl titanate, used in this study were the same as those reported previous1y.l Commercial reagent-grade iso- propyl titanate was also used as a starting complex.Doping experiments were performed by the introduction of the vapour of aluminium isopropoxide, or iron or chromium acetylacetonate. The former two complexes are commercial samples, and the latter was prepared according to a literature m e t h ~ d . ~ The dopant complexes were transported by nitrogen gas into the furnace and mixed with ethyl titanate vapour at the entrance to the reaction furnace. The substrate, consisting of a glass sheet (Corning no. 7059) and a titanium plate, was treated in the same way as in the previous study.' The electrical and electrochemical properties of the TiO, films were examined by the same method as described previous1y.l The electrolyte used for the electrochemical measurement was 2563 83-22564 ELECTROCHEMICAL PROPERTIES OF TiO, FILMS 0.1 mol dm-3 Na,S04 solution, unless otherwise specified.The reference electrode was a saturated calomel electrode (SCE). All potentials are given in V us. SCE. A 500 W Xe lamp and a fluorescent lamp were used as a light source for the photoelectrochemical study and photoconductivity measurements, respectively. Monochromatic light was obtained using a monochromator (Nikon G-250) from the 500 W Xe lamp. The gases used were purified by passing them through active copper and phosphorous pentoxide. RESULTS AND DISCUSSION ELECTRICAL PROPERTIES OF TiO, FILMS ON THE GLASS The electrical properties of TiO, films deposited on a glass substrate from ethyl titanate depend on the deposition conditions, especially the substrate temperature. Some examples of the data obtained are listed in table 1.The deposits obtained at 400-440 OC (hereafter referred to as lower temperatures) have a larger conductivity than those at 470-5OO0C (higher temperatures) while the conductivity of the latter TABLE 1 .-DEPOSITION CONDITIONS AND RESISTIVITY OF TiO, FILMS sample number 96 103 72 80 substrate temp./OC N, flow rate/cm3 s-l 0, flow rate/cm3 s-l film thickness/pm rutile content (%) resistivity/Q cm in H, in N, in Ar in 0, 420 0.60 0.30 0.80 62 2.5 x 103 (1.7 x 103)~ 2.8 x 103 (2.2 x 103) 1 . 9 ~ 103 3.3 x 103 (2.5 x 103) (6.2 x 10,) ~~ 500 0.60 0.30 0.90 0 2.6 x 105 3.6 x 105 5.3 x 105 (2.5 x 10) (4.6 x 10) (2.6 x 10) 1.5 x lo6 (4.0 x 10,) ~ 470 0.80 0.20 1 .oo 0 4.0 x los (1.3 x 10,) 4.4 x los 6.2 x lo6 6.6 x lo6 (2.2 x 103) (1.1 x 102) (1.3 x 104) 470 1 .oo 0 0.79 1 4.6 6.6 6.6 7.6 (3.3) (3.6) (3.3) (6.6) Resistivity values in parentheses are those under illumination.increases by lo4 under illumination even by a weak fluorescent light (light intensity, 0.6 mW crn-,). Exposure of the films to atmospheric gases also affects the conductivity. This effect is more notable under illumination than in the dark. The fact that the reducing gas, H,, increased the conductivity (especially under illumination) is consistent with the properties of an n-type semiconductor.1° The films deposited at lower temperatures had little photoactivity, probably indicating the existence of donors at such shallow levels as to be excited into the conduction band by thermal energy at room temperature. The films grown in the absence of additional oxygen gas have larger electrical conductivity owing to a drastic deviation of the stoichiometry, which was confirmed by a weight increase when oxidized in air, and had very little photoresponse.The TiO, films obtained from isopropyl titanate tended to have a very low conductivity (lO-ll C2-l crn-l).l1 The photoresponse of the high-temperature deposits is large, as described above, but the response rate was unexpectedly slow. A typical example of theY. TAKAHASHI et al. 2565 1 o9 1 G lo8 3 5 4 d m m b) .- 1 o7 1 o6 Jlight - On /-- I !' O2 I I 1 I- - 1 h time FIG. I.-Photoresponse in the electrical resistance of a TiO, film (no. 103) under an atmosphere of H,, N, and 0, gases (for deposition conditions see table 1). time dependence of the photoresponse in various gases is shown in fig.1. When the sample was illuminated, the resistivity decreased sharply in the initial stages and the rate of decrease became slower. 2-3 h later the resistivity levelled to a minimum value that was dependent on the type of atmospheric gas used. On the other hand, when the light was turned off, the change in resistivity was much slower, especially under hydrogen, nitrogen or argon. The resistivity measured after the d.c. ohmmeter was detached for a short while or when the polarity was reversed was still consistent with that assumed by extrapolating the transient curves or that obtained before changing the polarity, suggesting strongly that the charge carrier is not ionic but electronic. The very slow response on turning off the light tells us that the carriers which were generated under illumination have a long lifetime even in the dark.This process cannot be explained by a simple physical mechanism, but rather in terms of a chemical surface reaction of the films with the atmospheric gas (an adsorption-desorption process). Hydrogen is thought to diffuse into the TiO, crystal, donating an electron by ionization in the lattice.l0'l2 There have been no reports of such behaviour for nitrogen: traces of hydrogen in the gas may play a role. In contrast to the case of hydrogen or nitrogen, the response in oxygen is low but very fast, as shown in fig. 1, indicating that the reaction between the film surface and the atmospheric oxygen can be rapidly equilibrated, including deprotonation.12 Doping with aluminium, chromium or iron had a slight effect on the dark- and photo-conductivities, but the essential features of both conductivities were the same as for undoped films, as shown in table 2.Note, however, that the response rate when the light was turned on and off is enhanced by the aluminium doping, as shown in fig. 2. We have no clear explanation of this doping effect at present, but more detailed studies may enable us to find a suitable composition of the films for them to work effectively as a gas sensor.2566 ELECTROCHEMICAL PROPERTIES OF TiO, FILMS TABLE 2.-DEPOSITION CONDITIONS AND RESISTIVITY OF DOPED TiO, FILMS sample number 216 234 253 293 294 dopant substrate temp./”C N2 flow rate/cm3 s-l film thicknesslpm rutile content (%) resistivity/S2 cm in H, in N, in Ar in 0, A1 A1 435 480 0.6 0.6 1.3 1.4 100 0 2.0 x 103 1.9 x 106 5.0 x 103 2.4 x 107 5 .4 ~ 103 2.4x 106 4.5 x 103 2.4 x 107 (1.5 x (1.3 x lo2) (2.0 x 10,) (9.1 x 10,) (1.8 x lo2) (3.1 x lo2) (7.2 x lo2) (4.1 x lo4) Cr 43 5 0.6 0.86 100 1 . 7 ~ lo2 4.0 x lo2 (1.3 x lo2) 3.1 x lo2 4.3 x 102 (2.9 x 10,) (9.9 x 10) (1.1 x 102) Fe Fe 415 470 1.1 1.1 3.9 3.3 0 0 1.7 x 103 1.7 x 104 7.5 x 103 4 . 4 ~ 104 (1.1 x 103) (1.5 x 102) 1 . 2 ~ 103 1.1 x 105 (1.4 x lo2) (3.6 x lo2) 6.6 x lo3 (1.1 x 10) (1.7 x 10) 1.8 x lo4 (1.7 x lo2) (1.5 x lo2) a Resistivity values in parentheses are those under illumination. lo1* 1o1O Y $ 109 1 o8 1 o7 c: --.. .r( v, 2 1 o6 0 1 2 3 time/h FIG. 2.-Photoresponse in the electrical resistance of an Al-doped TiO, film (no.234) under an atmosphere of H,, N,, Ar and 0, gases (for deposition conditions see table 2).Y. TAKAHASHI et al. 2567 I 3 50 400 wavelength/nm FIG. 3.-Spectral dependence of the photoconductivity of a TiO, film (no. 72) (for deposition conditions see table 1). Data were not corrected for light intensity. 100 - ---/-- 60 '. ,- 0 0 - 1 I I I I 300 400 500 600 700 800 wavelength/nm FIG. 4.-Visible absorption spectra of TiO, films: (-) no. 72, (---) no. 94, (--. .-) no. 96, (-*--) no. 103, (. * ' - -) no. 139. Deposition conditions: no. 72, 96 and 103, see table 1 ; no. 94, substrate temperature 390 O C ; no. 139, substrate temperature 420 O C , from isopropyl titanate. The spectral dependence of photoconductivity measured in air using the Xe lamp as a light source is shown in fig.3. Maximum response is observed at 350 nm, namely at the intrinsic region as expected. However, the photoconductivities shown in tables 1 and 2 are those obtained under illumination by white light from a fluorescent lamp. Therefore, some deeper extrinsic levels might also contribute to the photoconductivity. l3 The difference between low-temperature deposits and those at high temperature is clearly shown in the visible-light absorption spectra (fig. 4). The deposits at higher temperatures showed a clear absorption edge around 400 nm, which is consistent with the photoconductivity data in fig. 3. The low-temperature deposits had no clear absorption edge, indicating that there are many trap levels in the forbidden region.The temperature dependence of the resistivity (20-200 "C) of the films obtained at higher temperatures suggested the band gap to be 3.0-3.3 eV. For the low-temperature2568 ELECTROCHEMICAL PROPERTIES OF TiO, FILMS deposits on the other hand, the activation energy of conduction is as low as 0.3 eV, indicating the existence of shallow traps. ELECTROCHEMICAL PROPERTIES OF TiO, FILMS ON TITANIUM PLATE As reported previously,' the TiO, films obtained with our method behave as typical n-type semiconductor electrodes in a photoelectrochemical cell. The characteristics of the anodic current against potential were found to depend strongly on the deposition conditions, in particular on the deposition temperature. The general features observed in 0.1 mol dm-3 Na,SO, solution can be summarized as follows: (1) The films deposited at lower temperatures exhibit large anodic photocurrents which are saturated at a rather positive potential (ca.0.5 V us. SCE). (2) The films deposited at higher temperatures showed rather small photocurrents which are saturated at a considerably negative potential (ca. 0 V us. SCE). (3) Anodic dark currents are negligibly small in every case. (4) The deposits are stable against photocorrosion. ( 5 ) The deposits at an optimum condition have a high quantum efficiency for anodic photocurrent, comparable to that of single crystals. 100 80 h E 5 60 E E U .d 40 m u. 20 0 310 330 350 370 390 410 wavelength/nm FIG. 5.-Spectral dependence of the quantum efficiency for the anodic photocurrents of TiO, film electrodes: 0, no.Ti-04; 0, no. Ti-10; 0, no. Ti-12; ., no. 169. The anodic photocurrents were measured at 0.8 V us. SCE in 0.1 mol dme3 Na,SO, solution. Deposition conditions: no. Ti-04, Ti-10 and Ti-12, see ref. (1); no. 169, substrate temperature 470 OC, from isopropyl titanate. Thus these results indicate that the as-grown films have excellent properties as photoanodes. The decomposition temperature also affected the spectral dependence of the quantum efficiency for anodic photocurrents, as shown in fig. 5. Raising the deposition temperature leads to an improvement in the response at longer wavelengths.' A pure anatase film obtained at 420 OC from isopropyl titanatel' had a maximum response at ca. 365 nm, which is very close to that of the deposits at high temperaturesY.TAKAHASHI et al. 2569 from ethyl titanate. The film showed a very low conductivity as described above. Therefore it may be suggested that the spectral dependence of the photocurrent is closely related to the electrical properties of the films. An estimation of the donor concentration in the deposits may clarify the relation. A Mott-Shottky plot is one possible way to obtain such an estimation. We then measured the differential capacitances of these film electrodes in an aqueous electrolyte in the usual manner. However, we could not obtain a meaningful result, since the capacitance ( C ) measured depended strongly on the frequency (0.1-10 kHz) of the signal used in the impedance bridge, and the plot of 1/C2 against potential gave a straight line for one sample but not for another. 100 80 h E 5 60 E E 2 40 E 0 .- e r 20 0 31 wavelength/nm FIG.6.-Effect of Fe- and Al-doping on the action spectra for anodic photocurrents: 0, Fe-doped (no. 299); 0, Al-doped (no. 238); 0, undoped (no. 230). The anodic photocurrents were measured at 0.8 V us. SCE in 0.1 mol dm-3 Na,SO, solution. Substrate temperature: no. 230 and 238, 435 O C ; no. 299, 440 O C . The doping effects of aluminium and iron on the spectral dependence of the anodic photocurrent are shown in fig. 6, where only the results obtained at optimum doping conditions are shown. The dopant concentrations were too low (perhaps ca. 1% or so) to be measured precisely with our energy-dispersion type of EPMA. For reference the result on an undoped film, obtained under the same deposition conditions as the doped one, is also shown in fig.6. By doping with aluminium, the spectral response can be appreciably extended. However, on comparing fig. 6 with the result in fig. 5 , it is clear that the spectral dependence of the aluminium-doped film is similar to that of an undoped film at a higher temperature (470 "C). Therefore, we conclude that aluminium doping is not essential to extend the spectral dependence, in agreement with the observations of other workers.2*5 On the other hand, iron doping has such a remarkable effect that a higher quantum efficiency (q) is observed even at longer2570 ELECTROCHEMICAL PROPERTIES OF TiO, FILMS wavelengths (25% at 400 nm and 0.7% at 500 nm). Interestingly this film is composed of 100 % anatase.Chromium doping has been reported to be effective in extending the spectral 5 7 However, in our case no such optimum doping conditions were found: doping with chromium suppressed the anodic photocurrents. These findings suggest that the doping conditions must be carefully examined, and that the doping effect may depend on properties of the original undoped samples. As mentioned above, the i against V characteritics of TiO, films can be controlled by the deposition temperature. A negative onset potential and sharp current rise on sweeping the potential are observed for films obtained at high temperatures. The film obtained from isopropyl titanate behaved similarly. If a film having such good i against V characteristics is used as a photoanode in a photovoltaic cell, the cell can be expected to have good current-voltage characteristics, namely a high fill factor.Therefore, we examined the output power characteristics of the cell, in which no. 169 film and platinum were used as the photoanode and photocathode, respectively. A single-compartment cell system was employed, with 0.1 mol dm-3 Na,SO, or TABLE 3 .--OUTPUT POWER CHARACTERISTICS OF ELECTROCHEMICAL PHOTOCELL, TiO, (NO. 1 ~~)~ELECTROLYTEIP~ electrolyte 0.1 mol dm-3 0.1 mol dm-3 0.1 mol dm-3 Na,SO, Na3P0, NaOH open-circuit voltage/V 0.44 0.81 0.78 short-circuit current/mA cm-, 1.02 1.01 0.96 maximum output power/mW cm-2 0.089 0.46 0.50 fill factor 0.20 0.56 0.67 @(%)a 0.089 0.46 0.50 a Maximum power conversion efficiency. Light intensity, 100 mW ern-,.0.1 mol dm-3 NaOH solution without intentional addition of any redox couple was used as an electrolyte. The results obtained are shown in table 3. The cell, TiO,(no. 169)lO.l mol dm-3 NaOHIPt has a very high open-circuit voltage, Kc, and a large fill factor of ca. 0.6-0.7. The light-electricity conversion efficiency (4) was 0.5 % under white-light illumination. Both Kc and q4 are comparable to the values (0.8 V and 0.5%) observed for a two-compartment cell system TiO,(rutile, single crysta1)lNaOH 11 H,SO,]Pt.l4 The high efficiency observed may be ascribed to the high fill factor as well as to the high value of Cc. Good i against Vcharacteristics of the film (no. 169) used contribute to the high fill factor. In order to clarify the unexpectedly high value of GC, the rest potentials of the TiO, and Pt electrodes were measured in the dark and under illumination.The data are listed in table 4. The potential of the TiO, electrode varied in the usual manner by ca. 60 mV per pH unit according to the Nernst equation,15 whereas that of the Pt electrode was less sensitive to any variation of pH. The difference between the pH dependences of both electrodes should be the reason why the high value of Kc wasY. TAKAHASHI et al. 257 1 TABLE 4.-REST POTENTIALS (V US. SCE) OF TiO, ELECTRODE (NO. 169) AND Pt ELECTRODE 0.1 mol dm-3 Na,SO, 0.1 mol dmP3 Na3P0, 0.1 mol dm-3 NaOH electrolyte (PH 7 ) (PH 12) (PH 12) ~~ ~ ~~ ~ ~~ ~ TiO, -0.30 (-0.54)a - 0.73 (- 0.93) - 0.70 (- 0.92) Pt f0.15 (+0.15) -0.15 (-0.15) - 0.07 (- 0.07) a Potential values in parentheses are those under illumination.observed. Table 4 also indicates that the cell has a large dark e.m.f. (ca. 0.45-0.55 V depending on the pH of the solution), which was not always reproducible. Therefore, the net V,, induced by photon energy may be estimated to be 200-250 mV, which was comparable to the value (ca. 200 mV) found for a polycrystalline Ti0,-based cell6 We cannot clarify at present the reason why there exists such high dark e.m.f., although it may be supposed that dissolved oxygen plays an important role. On the Pt electrode no gas evolution was observed, unlike the TiO, electrode, on which 0, gas evolution took place. The potential of the Pt electrode (-0.07 V us. SCE in 0.1 mol dm-3 NaOH, pH 12) seems to be far more positive than that of the H+/H, couple (ca.-0.95 V us. SCE, pH 12) and more negative than that of the O,/OH- couple (ca. 0.25 V us. SCE, pH 12). Therefore, the actual cathodic reaction mainly occurring on the Pt electrode may be concluded to be the reduction of 0, gas in solution. The fact that bubbling N, gas around the Pt electrode reduced the value of Ti0,(5 mol dm-3 NaOHIO,, Pt by ca. 0.2 V supports this conclusion. The report that the cell had confirm our observations. = 0.89 V,17 and other reports on the effect of oxygen gas,18* l9 are likely to 1 2 3 4 5 6 7 a 9 10 11 12 13 14 15 16 17 ia 18 Y. Takahashi, K. Tsuda, K. Sugiyama, H. Minoura, D. Makino and M. Tsuiki, J. Chem. SOC., Faraday Trans, I , 1981, 77, 1051. A. K. Ghosh and H. P. Marusuka, J. Electrochem. SOC., 1977, 124, 1516. H. P. Maruska and A. K. Ghosh, Solar Energy Mater., 1979, 1, 237. J. Augustynski, J. Hinden and C. Stalder, J. Electrochem. SOC., 1977, 124, 1063. C. Stalder and J. Augustynski, J. Electrochem. SOC., 1979, 126, 2007. A. Monnier and J. Augustynski, J. Electrochem. SOC., 1980, 127, 1576. Y. Matsumoto, J. Kurimoto, Y. Amagasaki and E. Sato, J . Electrochem. SOC., 1980, 127, 2148. Y. Matsumoto, J. Kurimoto, T. Shimizu and E. Sato, J. Electrochem. SOC., 1981, 128, 1040. T. Moeller, Inorg. Synth., 1957, 5, 131. R. Schumacher, Ber. Bunsenges. Phys. Chem., 1980, 84, 125. Y. Takahashi, S. Wakayama, A. Ogiso and K. Sugiyama, J . Chem. SOC., Faraday Trans. I , to be published. (a) D. S. Ginley and M. L. Knotek, J . Electrochem. SOC., 1979, 126, 2163; (b) A. Harris, J. Electrochem. SOC., 1980, 127, 2657. A. K. Ghosh, F. G. Wakim and R. R. Addiss Jr, Phys. Reu., 1969, 184,979. A. Fujishima, K. Kohayakawa and K. Honda, Bull. Chem. SOC. Jpn, 1975, 48, 1041. T. Watanabe, A. Fujishima and K. Honda, Chem. Lett., 1974, 897. R. S. Davidson, R. M. Slater and R. R. Meek, J . Chem. SOC., Faraday Trans. I , 1979, 75, 2507. D. Laser and A. J. Bard, J. Electrochem. SOC., 1976, 123, 1027. J. Keeney, D. H. Weinstone and G. M. Haas, Nature (London), 1975, 253, 719. W. Gissler, P. L. Lensi and S. Pizzni, J. Appl. Electrochem., 1976, 6, 9. (PAPER 1 / 1900)

ISSN:0300-9599    

DOI:10.1039/F19827802563

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. Soc., Faraduy Trans. I, 1982, 78, 2573-2582 Temperature-programmed Reaction Studies of the Interaction of Methyl Formate and Ethanol with Polycrystalline Zinc Oxide BY MICHAEL BOWKER, HILARY HOUGHTON AND KENNETH C . WAUGH* Corporate Laboratory, Imperial Chemical Industries PLC, P.O. Box No. 1 1 , The Heath, Runcorn, Cheshire WA7 4QE Received 10th December, 198 1 The reactive nature of the polycrystalline zinc oxide surface has been investigated using methyl formate and ethanol as probe molecules; these have shown the dominant influence of the cation-anion dual site. Methyl formate decomposes during adsorption at the dual site to adsorbed methoxy and formate species. On temperature programming these show identical kinetics and reaction pathways to the methoxy and formate species observed previously after the adsorption of formaldehyde or methanol.Adsorption of ethanol at the dual site results in the formation of an adsorbed ethoxy species. On temperature programming, at 510 K, the zinc oxide surface abstracts one of the /?-hydrogen atoms to form an unstable C,H,O fragment which decomposes mainly (90%) to ethylene; a small amount (cu. 10%) of the desorbing material which is either acetaldehyde or ethylene oxide (these cannot be distinguished by mass spectrometry) is desorbed coincidently at this temperature in an isomerization/desorption step. Despite the ethylene formation being a dehydration reaction, no water is observed in the desorption spectrum, the ethoxy oxygen species being left on the surface during decomposition.A previous paper1 reported the adsorptive interactions of carbon monoxide, carbon dioxide, hydrogen, formaldehyde and methanol on a highly defected polycrystalline zinc oxide. Saliently, both formaldehyde and methanol were observed to dehydrogenate upon adsorption on the surface. Particuarly significant was the observation that both molecules reacted with the surface oxygen of the zinc oxide to form an adsorbed formate; this formate decomposed at 580 K to yield mainly carbon monoxide and hydrogen, which desorbed coincidently. After methanol dosing, numerous methanol peaks were observed in the desorption spectrum in the temperature range 300-500 K, approximately one-tenth of this amount being desorbed coincidently as formaldehyde. For a formaldehyde dosed surface the situation was reversed, with a large amount of formaldehyde desorbing in the same temperature regime with coincident and smaller amounts of methano1.l The present paper further analyses the reactivity of the zinc oxide surfaces, using as a probe the adsorptive interactions of methyl formate and ethanol and employing the techniques of temperature-programmed desorption and temperature-programmed reaction spectroscopy.2$ The reaction pathways of methyl formate on an idealised oxide have been fully elucidated by Madix and co-workers,2 who have shown the necessity for nucleophilic attack of the oxygen anion as a prerequisite to further reaction.In the case of ethanol, Wachs and Madix3 have shown the promotional effect of the metal-cation-oxide-anion dual site for its chemisorption as an ethoxy species.The subsequent reactions of the chemisorbed ethoxy species have also been well characterised.2 This background serves as a uszful basis for an investigation of the nature of the chemisorption sites on polycrystalline zinc oxide. 25732574 INTERACTION OF METHYL FORMATE AND ETHANOL ON ZnO EXPERIMENTAL The equipment used has been described previously1 in detail. The temperature-programmed desorption and reaction experiments were all carried out at a heating rate of ca. 0.7 K 0, and all products were monitored on a Vacuum Generators model QX200 quadrupole mass spectrometer with good resolution. The ZnO sample was AnalaR (99.7% pure) and the major impurity (carbonate) is quoted at a level of 0.25%. The surface area of the sample was determined to be 3 f 0 .2 m2g-l by the B.E.T. method. The AnalaR sample was well defined, consisting of crystallites of typical length around 1000 to 3000 A and 1000 A diameter. Ref. (1) gives a more detailed description of this sample, including electron micrographs. The gases used in this study were all liquids at room temperature and were purified by cycles of freezing, thawing and pumping. The adsorbate was dosed onto the zinc oxide held at 295 K through a needle valve from the equilibrium vapour pressure above the liquid in a closed dosing line. Vapour purity could be checked by (i) constancy of the pressure above the liquid and (ii) analysis of the mass spectral cracking pattern in the adsorption system. Two types of blank expeiiment were conducted.The first tested the degree to which adsorption of the gases on the walls of the apparatus contributed to the desorption spectrum by dosing the total system, minus zinc oxide, as described above, evacuating the system and temperature programming. Nothing other than a rising background was detected. The second examined the possibility that the exposed Pt/Rh heating wire might contribute to the desorption spectrum by cracking the desorbing molecule. A realistic pressure of adsorbate, ca. Torr, (comparable to that measured during the desorption process) was maintained in the apparatus, the tube, which normally housed the zinc oxide, in this case containing the same weight of glass beads with the thermocouple embedded in the middle of them. On temperature programming considerable cracking of ethanol and methyl formate to CO and H, was observed at ca.380 K. This derived from the mode of temperature control used in these experiments, which allowed the heating wire initially to become red hot before control was imposed. Therefore, although H, and CO were observed in the desorption spectra at 380 K coincident with the desorption of formaldehyde, methanol and ethanol they have not been included in fig. 1 and 2 since they appear to be an artefact of the system. Above 400 K, the heating wire never became so hot as to cause significant product cracking. METHOD The method of catalyst pre-treatment employed here was the same as that reported previously.' The zinc oxide was heated in uacuo from ambient to 700 K at a rate of 0.7 K s-l.Carbon dioxide and water were the main molecules to desorb, having the same peak maximum temperatures as observed before.' Subsequently, after several heating and cooling cycles, the zinc oxide was dosed with hydrogen (ca. 1 Torr) at 550 K for 15 min, after which the gas-phase hydrogen was pumped away, any adsorbed hydroger, being removed by raising the temperature to 700 K in uacuo. As reported previously,' this pre-treatment produced a zinc oxide which was described as defected, in that after adsorbing hydrogen over a range of adsorption temperatures no water was seen in the desorption spectrum (that is, the hydrogen was unable to remove any further oxygen anions from the surface or the bulk). However, during the initial thermal treatment water, resulting from hydroxy ion combination, had desorbed at 460 and 570 K.RESULTS AND DISCUSSION ADSORPTION OF METHYL FORMATE The products which desorbed from the surface after dosing with methyl formate at 295 K are shown in fig. 1 . As with the desorption spectra of both formaldehyde and methanol,' the dominant feature of the methyl formate desorption spectrum occurs at 570 K, and corresponds to the decomposition of a surface formate species. The coincident desorption of carbon monoxide and hydrogen with smaller amountsM. BOWKER, H. HOUGHTON AND K. C. WAUGH 2575 CO, (M/e = 44) : x 10 CO (M/e = 28) f 3-3/ / ) CH30CH,(M/e = 45) X 10 - ?CO (M/e = 30) X 10 I I I 1 I I I 350 4 0 0 4 5 0 5 0 0 5 5 0 6 0 0 6 5 0 temperature/K FIG. 1 .-Desorption spectra resulting from the adsorption of methyl formate at room temperature.of water and carbon dioxide mark this species as formate. Methanol and formaldehyde also desorb over a broad range of temperature, as seen previously.' The formate species can be formed by two routes after adsorption of methyl formate : (i) nucleophilic attack of the adsorbed molecule by surface oxygen species leads to adsorbed methoxy and adsorbed formate species HCOOCH,(g)-+ HCOOCH,(a) HCOOCH,(a)+O(s)-+ HCOO(a)+CH,O(a) (ii) subsequent degradation of the adsorbed methoxy species by the surface anions, as described previously,1 also produces the formate species CH,O(a) -+ H(a) + H,CO(a) (3) H,CO(a) + O(s) -+ H,CO,(a) H,CO,(a) -+ H(a) + HCO,(a) (a) and (s) refer to adsorbed and lattice species at the surface, respectively.That the reactions (3)-(10) do occur can be seen by inspection of fig. 1 and table 1, where the amount of carbon monoxide evolved from the decomposition of the formate species at 570 K is approximately four times greater than the amount of2576 INTERACTION OF METHYL FORMATE A N D ETHANOL ON ZnO TABLE 1 .-ADSORBED SPECIES, COVERAGES, MAJOR DESORPTION/DECOMPOSITION PRODUCTS, PEAK MAXIMUM TEMPERATURES AND DESORPTION ACTIVATION ENERGIES major desorptiona adsorbed desorption coverage/ peak maximum activation adsorbate species products molecule cm-2 temperature/K energiesfkJ mol-I methyl methoxy H2C0 formate CH30H formate H2 co co2 H2O CH30H ethanol ethoxy C2H,0H C2H4 C2H4O H2 1 x 1013 2~ 1013 5 x 1013 1.3 x 1014 2.5 x 1013 I x 1013 I x 1014 I x 1013 5 x 1013 6 x 10l2 6 x 1013 340,380 570 7, 7 9 , 9 , 9 7 360,410 510 94, 105 159 9 , 77 7, >, 7, 100, 114 142 9 9 ?, a Calculated assuming a value of 1013 s-’ for the pre-exponential.methanol desorbing at 340 K, this being the dominant product of methoxy non- oxidative desorption/decomposition. On the basis of reaction (2), which represents the methyl formate to be decomposing on the cation-anion dual site with the methoxy species splitting off and adsorbing on the zinc cation, the methoxy:formate ratio should be 1. Even when allowance is made for the methanol and formaldehyde which desorb over the temperature range 400-600 K, the formate species still exceeds the methoxy species by a factor of 2. A complete listing of coverages and desorption activation energies is given in table 1.The coverages listed are mean values of ten experiments of the same dose, any one coverage varying by +20%. The inaccuracy involved here derived from the irreproducibility inherent in our dosing procedure. HCO,(a)-+ H(a)+CO,(a)+V,(s) (6) HCO,(a) --+ H(a) + CO(a)+ O(s) (7) (8) CO(a) -+ CO(g) (9) Coda) -+ CO&) (10) H(a) + H(4 -+ H,(g) where V,(s) is a surface oxygen vacancy. Tamaru and co-workers4 have shown that the selectivity to carbon dioxide in formate decomposition is a function of formate coverage, but as implied in reaction (6) it appears also to be a function of the degree to which the zinc oxide surface is defective. The assignment of the 570 K peak as formate in the desorption spectrum of methanol, formaldehyde and carbon dioxide/hydrogen co-adsorption1 is confirmed in this work.Indeed, the desorption of hydrogen and carbon monoxide together with Iesser amounts of water and carbon dioxide at 570 K can be taken as a ‘finger-print’ for the existence of a formate on a polycrystalline zinc oxide surface.M. BOWKER, H. HOUGHTON A N D K . C. WAUGH 2577 At 340 and 380 K formaldehyde, methanol, water and hydrogen desorb, again exactly as had been observed previously with formaldehyde and methanol adsorption.' These peaks are characteristic of the decomposing and further reaction of an adsorbed methoxy species. Reactions (1 I)-( 14) have been outlined beforel but the sites responsible for them have not been defined: CH,O(a) -+ H(a) + H,CO(a) (1 1) H,CO(a) --+ H,CO(g) (12) (13) CH,OH(a)--+ CH,OH(g).(14) H(a) + CH,O(a) -+ CH,OH(a) An indication as to the nature of their identity is to be found by comparison of the decomposition products of the methoxy species resulting from methanol adsorption on partially oxidised copper, Cu( 1 and from methyl formate adsorption on partially oxidised silver, Ag( 1 lo)., The former showed the necessity of the presence of oxygen for the decompositional adsorption of methanol into surface methoxy and hydroxy species, the methoxy species ultimately decomposing/desorbing as formal- dehyde, hydrogen and methanol at 340 K5 (The absence of oxygen meant that while methanol adsorbed dissociatively into methoxy species and hydrogenatoms, desorption was recombinative, only methanol being produced.6) The silver work also showed the need for a sparsely populated, yet oxygenated, surface for the dissociative adsorption of methyl formate into surface methoxy and formate species.In this case, therefore, it is suggested that the surface is probably a sparsely oxygenated zinc, possibly the zinc-dominated polar face. Methyl formate adsorption on to this face will be described by reactions (1) and (2), but because of the lack of surface oxygen some of the methoxy species so formed will not be oxidised to the formate and decomposes/desorbs coincidently as formaldehyde and methanol. The reactivity of these sites is remarkably similar to the oxygen-dosed copper( 1 lo), both in terms of the products desorbed and in terms of their energetics (T, = 340 and 380 K for zinc oxide compared with T, = 340 K for the oxygen-dosed copper5). ADSORPTION OF ETHANOL The desorption spectrum obtained on temperature-programming the zinc oxide which had been dosed with ethanol (lo7 Langmuir) at 295 K is shown in fig.2. As observed in the methanol1 and methyl formate desorption-decomposition spectra, two temperature regimes of different reactivity are observed: the first at 380 and 420 K is identical to that seen with methanol, while the second at 5 10 K is considerably lower than observed with methanol. The mechanism of ethanol adsorption on partially oxidised metals [Cu( 110) and Ag(1 lo)] has been elucidated by Wachs and Madix., Using deuterated ethanol, C,H,OD, they have shown its adsorption to be dissociative, forming a surface ethoxy species and OD. It is not unreasonable to suggest therefore that the mechanism of ethanol adsorption on zinc oxide is the same as that on partially oxidised copper, namely that it is dissociative, forming a surface ethoxy species and OH, reaction (1 5): C,H50H(g) + O(s) * C,H,O(a) + OH(a).(1 5) The ethanol desorption peaks at 380 and 420 K (there may in fact be only one peak maximum at 390 K since the heating-rate variations at this temperature could be responsible for the 420 K peak) result from reversal of reaction (1 5 ) and the desorption2578 INTERACTION OF METHYL FORMATE AND ETHANOL ON ZnO FIG. 2.-Desorption spectra resulting from the adsorption of ethanol at room temperature. activation energies involved (94 and 105 kJ mol-l) are those for this recombination of the surface ethoxy and OH species.Since both ethanol and methanol desorb at this temperature the activation energy could be that for the migration of hydrogen back to the surface alkoxy species. However, only 40% of the total amount of the adsorbed ethanol is desorbed as the parent molecule at these low temperatures. At 510 K, the remaining ethoxy species decompose/desorb as hydrogen, ethylene and C,H,O, either ethylene oxide or acetaldehyde, which cannot be distinguished because of their near-identical cracking patterns. (Acetaldehyde appears to be a likely product since in its desorption spectrum ethylene is obtained at the identical temperature of 510 K deriving from the decomposition of the surface ethoxy specie^.^ However, in the absence of having obtained the desorption spectrum of ethylene oxide, no firm conclusion can be made. Indeed, that experimental evidence might yet be inconclusive, since ethylene oxide could rearrange on adsorption to acetaldehyde and thereafter to the ethoxy species.) While observing that ethylene and acetaldehyde desorb coincidently after ethanol adsorption on oxidised copper and silver single crystals, Wachs and Madix3 make no comment about the possible common origin of these products. The origin of the ethylene is ascribed to the dehydration of a surface C,H,OD, species, proposed to be formed by the interaction of two adsorbed C,H,OD molecules; acetaldehyde is thought to derive from the decomposition of the ethoxy species.The coincidence of the peak maximum temperature is the result of the two speciesM.BOWKER, H. HOUGHTON AND K . C. WAUGH 2579 having a common rate-determining step. What has to be explained in the present case is the common origin of ethylene, C,H,O and hydrogen at 5 10 K, the former two being in the ratio 10: I , bearing in mind that the surface species is an ethoxy species bonded, possibly, to the zinc-dominated polar face. Since hydrogen is also desorbed at 510 K it seems apparent that the rate-determining step must be carbon-hydrogen bond (C-H) scission of the ethoxy species. The preponderance of ethylene in the desorption spectrum coupled with the observation that on copper surfaces3 the scission of the a-C-H bond results in the formation of acetaldehyde, leads us to conclude that in this case on zinc oxide the rate-controlling step is, surprisingly, the breaking of the P-C-H bond of the ethoxy species. Were the reaction simply a thermally induced unimolecular decomposition, then the a-C-H bond, being the weaker (its bond strength is 368 kJ mol-l compared with 385 kJ mol-1 for the P-C--H8), would be expected to have been the one which broke.Also, since the desorption/decomposition activation energy, estimated from the peak maximum temperat~re,~ is only 142 kJ mol-l, one-third the bond dissociation energy, the transition state to decomposition must involve interaction of the P-H atom with the zinc oxide surface. FIG. 3.-Configuration of the adsorbed ethoxy species on the zinc polar face [the (0001) face] which would result in abstraction of hydrogen from the B-carbon atom by an oxygen anion in the surface.(The zinc ions are the lighter coloured spheres.) Fig. 3 is a scale ball-and-stick model of an ethoxy -species adsorbed on the zinc-dominated polar face of zinc oxide. The bond lengths in the ethoxy species are the sums of the covalent atomic radii of its constituent atoms. The molecule is depicted as bonded to a zinc ion (the lighter balls) on the polar face. The oxygen to zinc bond of the surface complex is taken to be at right angles to the plane of the polar face to maintain the tetrahedral co-ordination of the zinc ion, and the bond length is set to be the same as in zinc oxide itself, 2 A. The model clearly shows that the 8-hydrogen2580 INTERACTION OF METHYL FORMATE AND ETHANOL ON ZnO atom can interact, with no configurational constraints, with a neighbouring oxygen ion (the darker balls) and, indeed, the ethoxy species could well be stabilised by hydrogen bonding.Abstraction of this hydrogen atom by the oxygen anion is therefore both feasible and relatively facile, having an activation energy only + the 8-C-H bond dissociation energy. While configurations for both a- and 8-hydrogen abstraction could exist on the prism and polar faces, single crystal work has shown the former to be completely stable and not to be involved in significant reaction.1° Indeed it is only when the ethoxy species is adsorbed on the zinc cation of the zinc-dominated polar face that the 8-H atom is exclusively discriminated for in terms of its interaction with the oxygen anions of the surface (see fig. 3). For all other faces, the oxygen-dominated polar face and the prism faces, the a- and the 8-H atoms have an equal possibility of interacting with the oxygen anions of the surface. Therefore, for reasons which are not obvious but which may be associated with the possibility of there being accessible 0- sites in the surface,ll the active surface appears to be the defected (highly reduced) zinc-dominated polar face.0 FIG. 4.-Configuration of the adsorbed ethoxy species on the zinc polar face [the (0001) face] which allows the closest approach of the hydrogen atom of the a-carbon atom to any oxygen ion on the surface. The separation of the hydrogen atom from the oxygen anion is 1.5 A. Fig. 4 depicts the closest distance of approach of the a-hydrogen atoms to any oxygen anion on this polar face. It is 1.5 A from the anion.Although the rocking vibrational mode of the ethoxy species, as a whole, about the zinc-oxygen bond will reduce this distance, the probability of a-hydrogen atom abstraction will remain considerably lower than that of P-hydrogen atom abstraction. At 510 K the P-H atom is abstracted and hydrogen is desorbed. An unstable C,H,O species is thus left on the surface which apparently has two channels to its decomposition : (i) carbon-oxygen bond scission resulting in ethylene desorption andM. BOWKER, H. HOUGHTON A N D K. C. WAUGH 258 1 (ii) zinc-oxygen bond scission resulting in the desorption of C,H40, either acetaldehyde or ethylene oxide. From the relative amounts of ethylene and C2H40 formed at 5 10 K (1 0 : 1, ethylene : C2H40), the activation energy difference between the C-0 bond scission and Zn-0 bond breakage can be estimated; it is found to be ca.10 kJ mol-l. This result is in good accord with published bond strengths;12 the C-0 bond strength in ethanol is listed as 376 kJ mol-l, the Zn-0 bond strength being 9 kJ mol-l greater than this, 385 kJ mol-l. Reaction (1 5) plus reactions (1 6) and (1 7) are sufficient to describe the evolution of all fragments at 510 K: CH,CH,O(a)+O(s)-+ CH,CH,O(a)+OH(a) While the production of ethylene might be termed a dehydration reaction, hy- drogen, not water, was seen in the desorption spectrum, bearing testimony to the high level of reduction of zinc oxide. Indeed, for every ethylene molecule produced, one oxygen atom is left on the surface [reaction (1 7)], so that from the listed amount of ethylene desorbed, 1014 atoms cm-, of oxygen must have been left on the surface.Since the polar faces of this AnalaR zinc oxide constitute between 10 and 30% of the total surface area [that is, 3 x 1014 atoms crn-,, see plate 1 of ref. (l)] and since the zinc-dominated polar face is half of this and the maximum oxygen coverage is half monolayer, then in the extreme case of a completely defected zinc polar face (that is, every oxygen of the sublayer having been removed), the maximum oxygen incorporation to reconstitute that surface would be 8 x 1013 atoms crn-,, i.e. 3 x 1014/4 atoms crn-,). While this is close to the number of oxygen atoms observed to have been deposited, it seems unlikely that the whole sublayer would have been abstracted by the hydrogen pretreatment.It is more probable that the oxygen atoms are redistributed over the entire zinc oxide surface by migration13 and that the average defect concentration is 10% of the surface. Finally, note that no carbon oxides nor formaldehyde were seen in the desorption spectra, showing that no carbon-arbon bond scission had occurred. Parrot et ai.l4 have propounded the view that metal oxides are generally dehydro- genating when the metal ion can readily change its valence state; this is in contrast to their supposed dehydrating character when the outermost orbitals of the metal ion are completely filled. This concept, though crude, is supported by this work. Ethylene, the result of a dehydration reaction, is the predominant product observed in this work, the zinc ion outer orbital being completely filled.However, no water was produced due to the highly defected nature of the sample at the outset; under steady-state conditions in a continuous-flow apparatus, the hydrogen product would further react to produce water tomaintain an invariant defect population. However, the predominant influence in oxides is that of the cation-anion dual site effecting dehydrogenation of the alcohol. The state of reduction of the oxide and its ability to be reduced further by hydrogen determine the ultimate fate of the oxygen removed from the alcohol, whether to form water or to re-oxidise the oxide. We thank Dr M. Parsley for many protracted and helpful discussions.2582 INTERACTION OF METHYL FORMATE AND ETHANOL ON ZnO M. Bowker, H. Houghton and K. C. Waugh, J. Chem. SOC., Furuhy Trans. 1, 1981,77, 3023. M. A. Barteau, M. Bowker and R. J. Madix, Surf. Sci., 1980,94, 303. I. E. Wachs and R. J. Madix, Appl. Surf. Sci., 1978, 1, 303, A. Ueno, T. Onishi and K. Tamaru, Trans. Furaduy SOC., 1970, 66, 756. I. E. Wachs and R. J. Madix, J. Catul., 1978, 53, 208. ti M. Bowker and R. J. Madix, Surf. Sci., 1980,%, 190. M. Bowker, H. Houghton and K. C. Waugh, to be submitted. V. I. Vedeneer, L. V. Gurvich, V. N. Kondrater, V. A. Frankevich and E. Y . Frankevich, Energy of Disruption of Chemical Bonds (Iz. Akad. Nauk SSSR, Moscow, 1962), p. 71. P. A. Redhead, Vacuum, 1962, 12, 203. lo M. Bowker and K. C. Waugh, unpublished results. l1 J. Cunningham, B. Doyle and E. M. Leahy, J. Chem. SOC., Furuday Trans. I , 1979, 75, 2000. l3 J. Cunningham, E. L. Gould and E. M. Leahy, J. Chem. SOC., Furaday Truns. 1, 1979, 75, 305. T. L. Cottrell, The Strengths of Chemical Bonds (Butterworths, London, 1958). S. L. Parrott, J. W. Rogers Jr and J. M. White, Appl. Surf. Sci., 1978, 1, 443. (PAPER 1/1914)

ISSN:0300-9599    

DOI:10.1039/F19827802573

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I , 1982, 78, 2583-2592 Stereochemistry in Metathesis of n-Alkenes using Heterogeneous Oxide Catalysts BY F. KAPTEIJN AND J. C. MOL* Institute for Chemical Technology, University of Amsterdam, Plantage Muidergracht 30, 10 18 TV Amsterdam, The Netherlands Received 30th December, 198 1 The selective formation of geometrical product isomers in the metathesis of both propene and pent-2-ene (cis and trans) was studied over several solid catalysts. trans- And cis-but-2-ene were formed in nearly equal amounts from propene, although there were some differences between Re,O,/y-Al,O,, WO,/SiO, and MoO,/SiO, catalysts. The initial translcis ratio, (tic),, for but-2-ene appeared to be independent of the reaction temperature and the propene pressure. Over Re,O,/y-Al,O,, cis-pent-2-ene gave ( t / ~ ) ~ values for but-2-ene (B) and hex-3-ene (H) of 0.35-0.40 and 0.95-1.1, respectively.Starting with trans-pent-2-ene a preference was observed for the formation of trans-products: (c/t),B x 0.30 and ( c / t ) , H M 0.12. The weight content of the active phase in the catalyst had no influence on the stereoselectivity, indicating that the active centres have a very distinct character. These results are explained using a stereochemical model based on a metallacyclobutane intermediate. The governing factors in the selectivity of product formation are the repulsive interactions between the alkyl substituents of the intermediate structure and the interactions of the substituents with the catalyst surface. In the catalytic metathesis of alkenesl new molecules are formed via the rupture 2 R’HC=CHR =ir: R’HC=CHR’+ RHC=CHR.(1) Three different types of metathesis reactions can be distinguished: (1) productive metathesis, yielding new products, (2) degenerate (or non-productive) metathesis, in which the exchange of alkylidene groups does not result in new products, and (3) cis-trans isomerisation. In the metathesis of asymmetric internal alkenes these three reactions occur simultaneously, while in the metathesis of alk-1-enes only the first two reactions occur. An interesting point is that the stereoselective formation of geometrical isomers in productive metathesis may give insight into the structure of the catalytic centres. In many cases, however, it has been observed that the translcis isomer ratio of the metathesis products attains its thermodynamic equilibrium value when the metathesis reaction is still far from equilibrium.Only at very short reaction times can the ratio of the formation of cis- and trans-products be established properly. In recent years stereoselectivity has mainly been investigated for homogeneous metathesis with tungsten- and molybdenum-based catalyst^.^-^ From these studies some general trends concerning the formation of geometrical product isomers can be derived :5 (a) for low-molecular-weight product alkenes (but-2-ene, pent-2-ene) there is a slight preference for the cis-isomer if the starting alkene is mainly cis and for the trans-isomer if the reactant is mainly trans, irrespective of the size of the alkyl groups in the starting alkene, (b) with increasing chain length and increasing bulkiness of the alkyl groups in the product molecule, the preference for the trans-isomer increases, regardless of whether the reactant molecule is trans or cis.The stereoselectivity and reformation of carbonsarbon double bonds : 25832584 STEREOCHEMISTRY IN METATHESIS OF ALKENES depends, however, not only on the structure of the product alkenes but also on the specific catalyst system used. Moreover, differences are observed in the stereoselectivity of the same catalyst precursor complex when using different co-catalysts or different activation methods. Until now, only one relevant study has been published on heterogeneously catalysed metathesis. Nakamura et aZ.6 reported an increasing preference for trans products with increasing chain length in the metathesis reactions of alk-1 -enes over a Re,O,/Al,O, catalyst at 545 K.In this paper we describe the results of a study of the selectivity of the metathesis of both propene and pent-2-ene (cis and trans) over several solid catalysts; the results are explained using a stereochemical model based on a metallacyclobutane reaction intermediate. EX P E R I M E NT A L5 MATERIALS Propene (polymerisation grade, Gerling, Holz and Co.) contained 0.1 % propane and 0.01 % ethane. Water and oxygen were removed before use by molecular sieves and a Cu/Al,O, catalyst. cis-Pent-2-ene and trans-pent-2-ene (Fluka, g.c.-pure) were freeze-distilled three times and stored with activated molecular sieves. This purification was performed as the trans-isomer contained 5 % high-boiling components.After purification the cis-pent-2-ene contained 0.44 % of the trans-isomer, and the trans-pent-2-ene contained 0.25 % cis-pent-2-ene. Catalysts were prepared by impregnation of y-Al,O, (Ketjen CK 300) or SiO, (Grace 1 13) particles (0.15-0.18 mm) with a solution of the ammonium salt of the transition metal in doubly distilled water. The water was slowly evaporated under reduced pressure and the particles dried overnight at 400 K. The catalyst was calcined in air for 4 h at 823 K and activated in situ in helium at the same temperature. The following catalysts, with different weight contents of the active phase, were prepared : Re,O,/y-Al,O,, WO,/y-A1,0,, MoO,/y-Al,O,, WO,/SiO, and MoO,/SiO,.REACTION SYSTEMS Continuous flow (propene feed) and pulse experiments (pent-2-ene feed) were carried out in a microreactor system using 0.02-1.0 g of catalyst, diluted with inert glass beads of the same size. The total reaction presure was kept at 0.35 MPa. Different partial propene pressures were obtained in the flow experiments by diluting with helium. In the pulse experiments amounts of 1 mm3 (liquid) pent-2-ene were injected into a stream of helium flowing through the reactor. The conversion of the reactant in both types of experiments was varied by changing the flow rate through the catalyst bed. The reaction temperature was measured by chromel-alumel thermo-elements placed just under and above the catalyst bed. ANALYSIS Product analysis was performed using a gas chromatograph equipped with FID (HP 5750 B).The product gases were separated at room temperature on a stainless-steel column packed with 30% bis[2-(2-methoxyethoxy)ethyl] ether on Chromosorb P (0.15-0.20 mm), using nitrogen as carrier gas. In the flow experiments samples were introduced using a Carle sampling valve. In the pulse experiments the reaction products were collected in a cold trap at 77 K. By rapidly heating this trap with water at 333 K the whole sample was analysed on-line. The peak areas were real-time processed using a PDP 11/10 computer system equipped with a sophisticated g.c.-signal integrating program.’F. KAPTEIJN AND J. C. MOL 2585 3 equilibrium 1 I t I I I 0 10 20 x (70) FIG. l.--trans/cis Ratio of but-2-ene as a function of the propene conversion (x) over Rez07/y-Alz0,. Reaction temperature 323 K; partial propene pressure 0.05 MPa; wt% Re,07: V, 3; A, 8.3; 0, 10.7; 4, 13; H, 18.2.FIG. 2.-Isomerisation of cis-but-2-ene into trans-but-2-ene over 10.7 wt% Re,O,/y-Al,O, (W) and over y-alumina (0) as a function of the space time (W/F). Reaction temperature 323 K; partial butene pressure 0.01 MPa. RESULTS METATHESIS OF PROPENE The ratio of formation of the two but-2-ene isomers by metathesis of propene was studied over all the catalysts mentioned in the experimental section. The initial translcis ratio of but-2-ene7 (t/c)o B, was obtained by plotting the experimental data for the t / c but-2-ene ratio against the propene conversion and extrapolating to 0% conversion. To obtain proper ( ~ / C ) ~ B values, low propene conversions had to be established since a rapid metathetical cis-trans isomerisation took place (fig. 1).At low temperatures the action of the support could be neglected, e.g. over a 10.7% Re,O,/y-Al,O, catalyst the cis-trans interconversion of but-2-ene was ca. 160 times faster than over the support alone (fig. 2). At higher temperatures alumina is an active double-bond isomerisation catalyst. * Because of this, substantial amounts of but- 1 -ene * This also applies to silicas containing aluminium impurity,8 e.g. Spherosil XOA 400 (100-lo00 ppm Al) and Ketjen silica (100O-lOO00 ppm Al). We used Grace 113 silica (< 10 ppm Al), which was sufficiently inert.2586 STEREOCHEMISTRY I N METATHESIS OF ALKENES mOJsI0,I 1.5 1.0 I FIG.3.--trans/cis Ratio of but-Zene as a function of the propene conversion for several catalysts. Equilibrium values of K,, at different reaction temperatures: 2.73 (323 K); 1.53 (583 K); 1.42 (673 K). Wt% MOO,: *, 3; 0 , 6 . Wt% WO,: *, 1; V, 3; .,6; ., 12. Wt% Re,O,: 0, 1; ., A, 3; +, 5.4; V, 8.3; 0, 10.7. were formed over WO,/y-Al,O, and MoOJy-Al,O, at reaction temperatures of 583-673 K.5 Even at low propene conversions the linear butenes were found to be in thermodynamic equilibrium.@ Therefore, only the results for Re,O,/y-Al,O,, WO,/SiO, and MoO,/SiO, could be used (fig. 3). For these catalysts (t/c),B was 0.80 0.05,0.95 & 0.02 and 1.15 & 0.10, respectively [the limits are standard errors (o), calculated by linear least-squares minimisation]. Fig.3 shows that these values are independent of the weight content of the active phase. The reaction temperature had no observable influence on (tlc), B for these catalysts in the temperature range studied (fig. 4). A variation of the partial propene pressure in the range 0.02-0.35 MPa also had no influence on (t/c),,B. METATHESIS OF PENT-2-ENE These experiments were carried out in a pulse reactor system; catalyst samples of 1.0 and 5.8 wt% Re,O,/y-Al,O, were used at 323 K. To avoid possible induction effects, we injected 5 mm3 of pent-2-ene before the measurements were started. After injection of this amount the steady-state conditions were obtained : but-2-ene and hex-3-ene were produced in equal amounts, and the total amount injected was equal to the amount collected. In preliminary experiments with trans-pent-2-ene, it was frequently observed that the catalyst turned pink after injection of the trans-pent-2-ene and was inactive for metathesis.After first injecting cis-pent-2-ene the catalyst did, however, exhibit activity for the metathesis of trans-pent-2-ene, although its conversionF. KAPTEIJN A N D J. C. MOL 0 -2 j, 2.0 - 1.5 - 1.0 - 2587 + + 5.8 % Re20,/A1203 +.' + * + - a I 0 I 4 x (%) I 8 FIG. 4.--trans/cis Ratio of but-2-ene as a function of the propene conversion at different reaction temperatures with different catalysts. 6% WO,/SiO,: A, 573; H, 673 K. 5.8% Re,O,(Al,O,: 4, 323; 0, 343; V, 373 K. was always 5-10 times lower than that of cis-pent-2-ene. These effects are also reported in the literature by Bilhou et aZ.l0 The catalyst was always activated with cis-pent-2-ene; from the preliminary experiments it appeared that this procedure had no influence on the results of the experiments with trans-pent-2-ene.It is also assumed that transient phenomena did not disturb our pulse experiments : with the Re,O,/y-Al,O, catalysts injection of pulses of propene gave the same results as in the continuous-flow experiments described above. Fig. 5 presents the results for cis- and trans-pent-2-ene. To take into account the effect of the geometrical isomerisation of pent-2-ene itself; the t / c (or c/t) product ratios are plotted against the t / c (or c / t ) ratio of pent-2-ene. The initial product isomer ratios are given in table 1 where x, denotes the conversion of pent-2-ene into but-2-ene and hex-3-ene, and xct (or xtc) is the conversion of cis-pent-2-ene (or trans-pent-2-ene) into its geometrical isomer.The xtc/x, or xct/x, ratio remained constant for both catalysts up to 20 % total pentene conversion (xCt + x,). The results can be summarised as follows : (a) The selectivity for the formation of cis-but-2-ene from cis-pent-2-ene is about the same as for trans-but-2-ene from trans-pent-2-ene9 whereas this selectivity is very different from that for the formation of cis- or trans-hex-3-ene. In the latter case there is a greater preference for the formation of the trans-isomer as compared with the but-2-enes. (b) The results obtained with the 1.0 and 5.8 wt% catalysts are the same, within experimental error. ( c ) The productive metathesis of pent-2-ene is2588 STEREOCHEMISTRY IN METATHESIS OF ALKENES 35 1 , ., f:; * , o j j $ . ; , oj#?., B m , 0 1 . m , . m;m ;;LB, , *8#@: 0 3 01 0 0 0 0 2 0 0 2 C 01 0 0 01 z A S B )I..'. 0 3 mmB .* ti2 01 0 a2 0 0 2 0 01 0 01 t f c , pent-2ene clt, pent-2-ene FIG. 5.-A, transleis product ratio for the metathesis of cis-pent-2-ene over Re,O,/y-Al,O,. B, cisltrans product ratio for the metathesis of trans-pent-2-ene over Re,O,/y-Al,O,. Top, 1 .O wt % Re,O,; bottom, 5.8 wt % Re,O,. Left, but-2-ene; right, hex-3-ene. Reaction temperature 323 K. TABLE I .-STEREOSELECTIVITY DATA FOR THE METATHESIS OF PENT-2-ENE OVER Re,O,/y-Al,O, weight content Re,O, feed 1 % 5.8 % W ) o B 0.35 & 0.05 0.4 +_ 0.1 (xct IX,) 0.75 0.77 (c/t)o B 0.31 0.05 0.30 & 0.03 ( C / O 0 H 0.12 f 0.04 0.1 3 rf: 0 .0 2 (Xtc/XIn) 0.35 0.35 (tlc)o H 0.95k0.10 1 . 1 k O . 1 cis-pent-2-ene trans-pent-2-ene faster than its geometrical isomerisation (a metathesis reaction too), and differs for the two pent-2-ene isomers by a factor two. DISCUSSION A possible mechanism for the metathesis reaction must, of course, explain the observed stereochemistry. At present, most literature reports are concerned with the carbene mechanism.l The key intermediate in the carbene mechanism is supposed to be a metallacyclobutane complex. The stereoselectivity of metathesis might be related to the structure of this intermediate. Metallacyclobutane compounds of the transition metals are at present few in number. In the literature only the syntheses of stable tungsta-,ll titania-12 and platina-cyclobutanes13 have been reported.The photolysis products of tungsten and molybdenum complexesll b t l4 suggest that this decomposition proceeds in a way analogous to that proposed for metathesis. Moreover, there are indications that the titania complexes react with neohex-1-ene according to a metathesis scheme.12 Therefore, the structure of these metallacyclo- butanes may be indicative of the structure of the cyclic intermediates in metathesis. Until now only the structural parameters of a series of platinacyclobutanes have beenF. KAPTEIJN AND J. C. MOL 2589 a A a B FIG. 6.-Structures of metallacyclobutane intermediate : A, flat ring; B, puckered ring. reported.13 These parameters suggest that the intermediate structure is not flat [fig.6(a)], but that the ring is bent across the Cl-C3 axis [fig. 6(b)]. The larger the ring substituents, the larger this dihedral angle (10-30°). It is conceivable, then, that the repulsive interactions between the ring substituents as well as the interactions of the substituents with the ligands on the transition metal (or catalyst surface) govern the stereoselectivity of the reaction. Steric repulsions in the metallacyclobutane transition state will cause an increase in activation energy, and lower the rate of its formation. Based on the puckered ring structure [fig. 6(b)], the four possible intermediate structures resulting from the reaction of either cis or trans isomers with a carbene species to yield new products CHR [eqn (211 ’ \CHR+ M = c m ’ + RHc=cHR (2) \ / M=CHR + R’HC=CHR C M CHR’ are depicted in fig.7 and 8, respectively. It can be seen that from structures I1 and I11 the isomeric structure of the reactant is retained in the product. The isomeric product composition is not disturbed by ring puckering: structure I converts into structure IV and structure I1 into structure 111. The trends in stereoselectivity can be explained on the basis of these structures and the substituent distances. The formation of low-molecular-weight products (R = CH,, C,H,) will be governed by the difference between the largest interactions, uiz. the interactions of the substituent in the e(quatoria1) or a(xia1) position of C1 and C3 with the catalyst surface. Here, the distance between the substituents in the e-position and the catalyst surface is smaller than between the substituents in the a-position and the catalyst surface.From fig. 7 and 8 it is obvious that structure I11 has the largest repulsion content and will be the least favoured intermediate. Apparently, structure I1 is sufficiently favoured over structures I and IV for the net result to be that cis yields mainly cis and trans yields mainly trans products (table 1). The differences observed in the (t/c),H values, starting with either cis- or trans- pent-2-ene, compared with (t/c)o B can be explained with increasing e-a interaction between substituents on C1 and C2. In this case, the preference of structure I1 over structures I and IV will be relatively smaller for cis- than for trans-reactants, so that2590 STEREOCHEMISTRY I N METATHESIS OF ALKENES I II lit IV l a z l 1”l l e z l IezI trans cis cis trans FIG.7.-Configurations of the metallacyclobutane intermediate for the productive metathesis of I la: I cis cis-alkenes. II 111 ‘”I I“$ trans trans IV lezl cis FIG. 8.--Configurations of the metallacyclobutane intermediate for the productive metathesis of trans-alkenes. the preference for trans increases with increasing size of R, as has also been observed in homogeneous metathesis. This explanation is supported by the fact that in both cases the preference for trans-hex-3-ene is a factor of 2.5 higher, pointing to a similar interaction. The stereoselectivity of the productive metathesis of alk-1 -enes can also be explained using this model. Referring to fig.7, in which R’ should be replaced by H, it is easily seen that starting with small alk-1-enes, the (t/c)o product ratio will be ca. 1 (cf. metathesis of propene), and gradually increase with increasing size of the R- substituentssy l5 due to the enhanced (1 a)-(2e) or (1 e)-(2a) interaction. Our explanation differs from models reported in the literature, which follow the proposal of Katz et aZ.* based on the interactions in 1,3-disubstituted cyclobutane, in the way that it incorporates the interaction with the transition-metal ligands or catalyst surface. This idea is supported by the observations of Bassetls that theF. KAPTEIJN AND J. C. MOL 259 1 selectivity of catalyst precursor complexes used in homogeneous metathesis increases when they are deposited on a support.Because of the longer M-C bonds and a wider Cl-C2-C3 bond angle in our intermediate compared with cyclobutane, the (1 a)-(3 a) substituent distance is 2-3 times as large as the distance between the substituent in the C1 (C3) position and the catalyst surface. Hence, the (la)-(3a) substituent interaction will hardly contribute to the stereoselectivity of the metathesis reaction. The difference between the ratios of the rates of cis-trans isomerisation and metathesis of pent-2-ene, starting with either cis- or trans-pent-2-ene, can be explained by thermodynamics. The rate constant for the cis to trans isomerisation is larger than that for the trans to cis reaction, since their ratio equals the thermodynamic equilibrium constant K,, (= 4.08 at 323 K). Thus it is expected that the initial rate of geometrical isomerisation starting with trans-pent-2-ene is lower than that starting with cis-pen t-2-ene.The absence of a temperature effect in the metathesis of propene is not surprising: the (tic),, B value can be considered as the ratio of the rate constants for the metathesis of propene into sole trans- or cis-but-2-ene. Since this ratio is nearly one, the activation energies for these two reactions are about the same and the temperature is expected to have no observable influence on this ratio, provided there are no other effects to be considered. Note that some investigators have suggested the involvement of a bimetallic centre as the active species in the metathesis reaction," but since the present model can be applied to mono- as well as to bi-metallic centres, no distinction can be made between these possibilities on the basis of the stereochemical data.Solid catalysts often appear to have heterogeneous reaction behaviour. Indeed, we have observed an exponential increase in the reaction rate of the metathesis of propene over Re,O,/y-Al,O, with increasing rhenium content.l* In view of the results of our stereochemical study, it is doubtful, however, as to whether this heterogeneity is due to active centres with different structures. The present study shows that with increasing content of the active phase the active sites still behave quite similarly with respect to the stereoselectivity of this reaction (fig. 3, table 1). This is in agreement with the conclusions of a kinetic study1% that the heterogeneous behaviour of the Re,O,/y-Al,O, catalyst is caused by the adsorption and desorption processes, and not by the rate-determining surface reaction(s). This work was supported by the Netherlands Foundation for Chemical Research (SON), with financial aid from the Netherlands Organization for the Advancement of Pure research (Z.W.O.).We thank Jan de Boer for experimental assistance. * (a) J. C. Mol and J. A. Moulijn, in Advances in Catalysis, ed. D. D. Eley, H. Pines and P. B. Weisz (Academic Press, New York, 1975), vol. 25, p. 131 ; (b) J. J. Rooney and A. Stewart, in Catalysis, ed. C. Kemball (The Chemical Society, London, 1977), vol. 1, p. 277; (c) R. H. Grubbs, Prog. Inorg. Chem., 1978, 24, 1 . * (a) N. Calderon, J. P. Lawrence and E.A. Ofstead, Adu. Organomet. Chem., 1979, 17, 449 and references therein; (b) E. A. Ofstead, J. P. Lawrence, M. L. Senyek and N. Calderon, J. Mol. Catal., 1980, 8, 227. (a) M. Leconte and J. M. Basset, J. Am. Chem. SOC., 1979,101,7296; (b) M. Leconte and J. M. Basset, Ann. N.Y. Acad. Sci., 1980, 333, 165 and references therein; (c) J. M. Basset and M. Leconte, in Fundamental Research in Homogeneous Catalysis, ed. M. Tsutsui (Plenum Press, New York, 1979), p. 285. T. J. Katz, Adu. Organomet. Chem., 1977, 16, 283 and references therein. F. Kapteijn, Ph.D. Thesis (University of Amsterdam, 1980). R. Nakamura, K. Ichikawa and E. Echigoya, Nippon Kagaku Kaishi, 1978, 12, 1602. ' E. F. G. Woerlee and J. C. Mol, J . Chromatogr. Sci., 1980, 18, 258.2592 STEREOCHEMISTRY I N METATHESIS OF ALKENES A.J. van Roosmalen, M. C. G. Hartmann and J. C. Mol, J. Catal., 1980, 66, 112. E. F. Meyer and D. G. Stroz, J. Am. Chem. Soc., 1972,94, 6344. lo J. L. Bilhou, J. M. Basset, R. Mutin and W. F. Graydon, J. Am. Chem. SOC., 1977, 99, 4083. These authors suggest that trans-pent-2-ene cannot form active centres for metathesis, in contrast to cis-pent-2-ene. In view of the impurity of the trans-pent-2-ene used (see the Experimental section), contamination might also be responsible for these effects. The colour change of the catalyst indicates a reduction of the rhenium; this colour change was hardly observed after treatment with cis-pent-2-ene. Apparently, in the latter case the rhenium was protected against excessive reduction. Attempts to purify the trans-pent-2-ene by preparative gas chromatography failed, so no definite conclusions can be drawn on this point. l1 (a) M. Ephritikhine, M. L. H. Green and R. E. Mackenzie, J. Chem. SOC., Chem. Commun., 1976,619; (b) G. J. A. Adams, S. G. Davies, K. A. Ford, M. Ephritikhine, P. F. Todd and M. L. H. Green, J. Mol. Catal., 1980, 8, 15. l2 T. R. Howard, J. B. Lee and R. H. Grubbs, J. Am. Chem. Soc., 1980, 102, 6876. l3 (a) R. D. Gillard, J. Organomet. Chem., 1971, 33, 247; (b) J. A. MacGinnety, J. Organomet. Chem., 1973,59,429; (c) D. J. Yarrow, J. A. Ibers, M. Lenarda and M. Graziani, J. Organomet. Chem., 1974, 70, 133. (a) M. Ephritikhine, M. L. H. Green and R. E. Mackenzie, J. Chem. SOC., Chem. Cmmun., 1976, 926; (b) D. C. L. Perkins, R. J. Puddephat and C. F. H. Tipper, J. Organomet. Chem., 1980,186,419. l5 M. Leconte and J. M. Basset, Nouv. J. Chim., 1979, 3, 429. l6 J. M. Basset, Y. Ben Taarit, J. L. Bilhou, J. Bousquet, R. Mutin and A. Theolier, Proc. 6th Int. Congress on Catalysis (The Chemical Society, London, 1976), p. 570. l7 (a) F. Garnier, P. Krausz and J. E. Dubois, J. Organomet. Chem., 1979, 170, 195; (b) J. Levisalles, H. Rudler, Y. Jeannin and F. Dahan, J. Organomet. Chem., 1979, 178, C8 and 1980, 187, 233. F. Kapteijn, L. H. G. Bredt and J. C. Mol, Recl. Trav. Chim. Pays-Bas, 1977, 96, M139. Recently, similar results with a Re207/A1203 catalyst have been reported by R. Nakamura, F. Abe and E. Echigoya, Chem. Lett., 1981, 51. l9 F. Kapteijn, L. H. G. Bredt, E. Homburg and J. C. Mol, Ind. Eng. Chem., Prod. Res. Dev., 1981,20, 457. (PAPER 1 / 1998)

ISSN:0300-9599    

DOI:10.1039/F19827802583

出版商:RSC    

年代:1982

数据来源: RSC