摘要:

Gas Phase Reaction between Iodine and TetramethylsilanetPart 2.-Kinetics and the Bond Dissociation Energy D(Me,SiCH,-H)BY ALAN M. DONCASTER AND ROBIN WALSH*Department of Chemistry, The University of Reading,Whiteknights, Reading, RG6 2ADReceived 22nd April, 1976The kinetics of the title reaction, which leads to the formation of small quantities of trimethyl-silymethyl iodide and HI at equilibrium, have been studied in the temperature range 609-649 K.From absorbance measurements on I2 the approach to equilibrium was found to be characterisedby the rate law- 5' = k[I,l%[Me4Si]{1 - [I\lle3SiCH2111HIldt KILI[Me&Iwhere K is the experimentally determined equilibrium constant. This expression is consistent withan iodine atom abstraction mechanism and for the step11.+ Me4Si --f Me3SiCH2 + H1log(k,/dm3 mol-' s-') = (11.82kO.30)-(122.2+ 3.6 kJ mol-l)/RTln 10 has been deduced. Fromthis the bond dissociation energy D(Me3SiCH2-H) = 415 k 5 kJ molPl(99.2 kcal mol-') is obtained.A more limited study of the reaction between I2 and neopentane indicates that Dc-~(Me4c) - Dc-H(Me&) = 1.8 ? 0.7 kJ mol-'.Some thermochemical implications of the absence of alternative reaction pathways are pointedout.Equilibrium measurements in the reaction of iodine with tetramethyl silane aredescribed in the previous paper.' We report here kinetic studies of this systemdesigned in particular to evaluate the bond dissociation energy D (Me3SiCH2-H).The technique of reaction with iodine is well established 2-4 as a sound and reliablemethod for such a determination.Apart from its intrinsic interest, the value of D (Me3SiCH2-H) will indicate themagnitude of any stabilising interaction (such as p z - dz delocalisation 5, betweenthe odd electron centred on carbon and the adjacent silicon atom in the trimethyl-silylmethyl radical Me3SicH2.Such an interaction manifests itself as a reductionin the value of D(Me3SiCH2-H) compared with D(Me3CCH,-H). The existingevidence for such a reduction comes from radical abstraction studies where, for theradicals CH3,6-9 CD3 6* lo and CF3,6* *. 11, l 2 slightly faster gas phase reactions areobserved with tetramethylsilane than with neopentane. Activation energy differencesare small but have been correlated, by means of Polanyi plots,l with a bond dissocia-tion energy difference of 4-8 kJ mo1-1.6-12 Similar small reactivity differences areobserved in solution abstractions by t-Bu0*.l49 l5 The currently accepted value l6for D(Me,CCH,-H) is 41 5 4 kJ mol-1 and so the above evidence points to 409 & 6kJ mol-' for D(Me,SiCH,-H).i- The authors regret that no reprints are available.290A .M . DONCASTER AND R. WALSH 2909EXPERIMENTALThis is as previously described.lRESULTSKINETIC MEASUREMENTSWe have shown in the previous paper that the reaction of iodine with tetramethylsilane leads to the following equilibrium at low conversionsIn accordance with general findings in gas-phase iodination systems 2-4 a three halvesorder rate expression was adopted as a working hypothesis, in this case with allowancefor the reverse reaction, viz.I2 + Me4Si Me,SiCH,I + HI.3 [Me,SiCH,I] [HI] -- dCrzl - - k2[12]*[Me4Si]{l -dt KC121 [Me&]where k$ is the rate constant and K the equilibrium constant.Integration of thisequation yieldswhere fi and f2 are functions of the instantaneous iodine concentration, [t2It, andc is a constant dependent on initial and equilibrium reactant concentrations (seeAppendix for full details). These expressions were tested in two ways. In the firstmethod the 1.h.s. of the integrated equation was plotted against t and linear fits wereobtained up to within -70 % of equilibrium. These plots showed scatter usuallyamounting to - 3 % in the slope of the least squares line, but there were no systema-tic trends.However the value of the slope did depend fairly heavily on the equili-brium concentration of 12. Because of uncertainty in the latter a more realisticuncertainty was 10 %. Other methods of data handling were found to be eithermore cumbersome or less precise or both. In the second test the invariance of k+with different starting pressures of reactants was established. Table 1 shows thevalues of k,, derived from the slopes of the integrated plots. The values at any giventemperature show a reasonable consistency although the range of initial pressureswas limited by the necessity of obtaining adequate conversions ; thus [Me,Si],/[I,],was only varied from - 11 to 150. Within this limitation the figures in table 1 offergood support for the assumed rate equation. Several runs were performed in thepacked vessel (S/V 4 times greater than the unpacked vessel) at T - 610 K, thelowest temperature of the range, and it can be seen from table 1 that no significantincrease in k, results (although the packed vessel values average 7 % higher than theunpacked ones, this is well within the experimental uncertainty).The reaction there-fore shows no serious surface effects and can safely be taken to be homogeneous at,and above, this temperature. The effect of added HI was briefly investigated at639.4 K. The results are shown in table 2. Because of the suppression of equilibriumby HI, changes are smaller and the precision of results is reduced. This can be seenin the increased scatter in the rate constant, k3..Apart from its effect on the equili-brium, added HI should inhibit the rate. This is not obvious from the data but if therate constant is corrected (see discussion of mechanism) according towhere ([HI]/[12]),v is the average value of the ratio during each run, the k i valuesshown in the last column of the table are obtained. These scatter less although they1-92910 GAS PHASE REACTION OF I,+Me,SiTABLE 1 .-RATE DATA IN THE IODINE-TETRAMETHYLSILANE REACTIONreactant pressurestemp/K a [Izlo/Torr [Me&]o/Torr 106k3/2/T0rr-~/~ s-1610.0610.0609.8609.8610.2619.4619.8619.8619.6619.1629.7629.3629.3629.3629.5629.7639.7639.5639.6639.5639.3639.4639.7649.8649.8650.3650.0650.5650.14.895.335.525.2314.035.395.035.535.615.195.235.195.024.885.6813.992.465.365.285.135.334.8513.501.405.175.134.445.0413.87610.3 (P) 5.40610.3 (P) 6.4061 1.9 (P) 8.6961 1.5 (P) 8.14610.4 (P) 10.2872.1112.2170.5310.8211.589.193.5149.5210.4226.366.395.3169.5202.6238.3187.0118.061.1102.2122.9181.2243.2181.5200.0107.3133.4196.6294.0201 .o96.6282.654.8149.9263.72.632.442.382.372.024.264.794.374.314.698.757.708.158.308.367.5216.015.715.214.915.714.212.626.125.824.824.424.222.12.822.502.582.672.17a (P) indicates packed vessel run.These data were not included in the Arrhenius plot.TABLE 2 .T H E EFFECT OF ADDED HI ON THE RATE CONSTANT AT 639 KtIJ~/Torr [Me&] o/Torr [HI] 0 /Ton ~ O ~ ~ ~ / Z / T O I T - ~ / * S-1 1 0 6 k ~ / ~ ' / T O ~ - ' / ~ S-'6.71 311 1.14 17.5 18.46.20 308 4.19 16.3 18.16.66 300 8.32 14.9 17.62.48 233 3.70 14.9 18.22.65 291 4.67 12.7 16.A. M. DONCASTER AND R . WALSH 291 1are now slightly higher than in non-inhibited runs. Thus the effect of added HI isvery small and what trend there is points to a low inhibition constant (< 0.12, the valueused above).The rate constants, k,, were fitted by a linear least squares procedure to theArrhenius equation, a plot of which is shown in fig. 1.-4.6nb -5.0-5.8This yieldsI I1.56 1.60 1.64lo3 KITFIG. 1.-Arrhenius plot for k3.log(k+/Torr-* s-l) = (10.94+0.30)- (193.3 k3.6 kJ mol-l)/RTln 10where the error limits are one standard deviation.Similar kinetic experiments on the reaction between I2 and neopentane, restrictedto a single temperature of 645 K, using the same technique and analysis as above,led tok+(645.0 K) = (1.32k0.10) x Torr-* s-l,the mean of six experiments. When compared with results for tetramethylsilane(calculated from the experimental rate equation) the following ratio is obtained at645 Kk3(Me4Si)/k3(Me4C) = 1.48 +_O.19.REACTION MECHANISM AND BOND DISSOCIATION ENERGIESThe observed kinetics parallel very closely those of the reactions of iodine withhydrocarbons.2 Thus the following iodine atom abstraction chain mechanism seemsmost likel2912 GAS PHASE REACTION OF 12+Me4SiIz(+M) $ 2I*(+M)12I-+ Me4Si Me3SicH2 + HI34Me3SiCH2 +I2 * Me3SiCH21 +I-.A stationary state treatment of this mechanism giveswhere K = klk3/k2k4.In practice k,[HI]/k,[IJ can usually be neglected bothbecause k2/k3 is certain to be small and [HI] < [I2] under experimental conditionsexcept in the inhibition experiments. The above expression then simplifies to theone which has been substantiated experimentally in the absence of added HI. Thusk+ can be identified with klK&. From the known values l7 of K t the following ex-pression for k, is obtainedlog(kl/dm3 mol-l s-l) = (11.82*0.30)-(122.2+3.6 kJ mol-l)/RTln 10.Typically alkyl radicals react with HI with low activation energies,, of the order of4 to 6 kJ mol-l, and so it is reasonable to suppose E2 = 5 f 4 kJ mol-l.HenceAH,",, (529 K) = El-E2 = 117.2f5.4 kJ mol-I. This may be corrected to 298 Kusing the approximate formula AH"(T) - AH"(298.15 K) = AC;(T-- 298.15 K) whereAC; = C,"(Me3SicH2) + Ci(H1) - C;(Me,Si) - C;(12), provided AC; can be obtainedas a function of temperature and averaged over the temperature range of interest.C;(HI) and C;(I2) are tabulated and C;(Me3SiCH,) - Ci(Me4Si) can readily beestimated by standard thermochemical means l8 and so this procedure was carriedout resulting in AH"(629 K) - AH"(298.15 K) = - 0.1 1 kJ mol-l, an almost negligiblecorrection. Thus AH"(298.15 K) = 117.3k5.4 kJ mol-l. Since AH,",2 = D(Me3-SiCH,-H)-D(H-I), then from the known value l7 of D(H-I) = 298 kJ mol-l,D(Me3SiCH2-H) = 41 5 k 5 kJ mol-I.A simpler treatment can be used to obtain the bond dissociation energy D(Me3-CCH2-H) if it is assumed that the reaction mechanism is identical.In that casek3( Me4Si) - kk,(Me4C) k5_ -5where k5 is the rate constant for the step I=+ Me4C -+ Me,CcH, +HI. Thus fromthe results kl/k5 = 1.48k0.19, at 645.0 K. If we assume a collision theory ratio(see Discussion) of A /A5 = 1.06 then E5 -El = 1.77 0.73 kJ mol-l. Since iodineatom abstraction steps such as (1) and (5) are highly endothermic it is also reasonableto assume D(Me3CCH2-H) - B(Me3SiCH2-H) = E5 -El. Thus the bond dis-sociation energy difference = 1.77k0.73 kJ mol-l and D(Me,CCH,-H) = 417 k 5kJ mol-' .DISCUSSIONThe kinetics of this reaction offer strong support for the proposed atomic chainmechanism. In particular the high A factor, 1011*82 dm3 mol-I s-l, is typical ofthose commonly found for iodine atom abstractions.If collision theory is usedto calculate this A factor, a value of 1011.39 dm3 mol-'s-l is obtained." Thesimilarity between these observed and calculated figures for A provides the justifica-tion for application of collision theory to calculate A 1 / A 5 . If collision theory iA. M . DONCASTER AND R . WALSH 2913approximately valid for I- + Me&, then it should apply well to relative A factors forthis reaction and I.+Me,C; it is not unreasonable to assume that most of theuncertainties in the absolute calculation will cancel. Thus the calculation l9 basedonly on mass and collision diameter differences (see Appendix) gives A1/A5 = 1.04.Even had a full temperature variation in the I,/Me,C study been carried out, it wouldhave been impossible to determine A5 with sufficient precision to establish this ratioreliably.Hence the relative method gives the difference in the two bond dissociationenergies Dc+(Me4C) - D,_,(Me,Si) = 1.8 kJ mol-' to a much greater precision thanthat to which either is known individually. It is because iodine atoms are much moreselective in their abstractions than are CH3 or CF, radicals that this difference mustbe regarded as more reliable than the 4-8 kJ mol-l deduced from Polanyi plots byother workers.6* l 2The absolute value for D(Me,CCH,-H) of 417 5 kJ mol-1 is in good agreementwith the currently accepted value l6 of 415+4 kJ mol-l.The value of D(Me,-SiCH,-H) of 41 5 + 5 kJ mol-1 although not previously determined, is in tolerableagreement with what could previously be inferred from D(Me,CCH,-H) and thebond dissociation energy difference. By combination of the results obtained here andthe enthalpy of reaction the previously undetermined dissociation energies D(Me,-SiCH2-I) and D(Me3CCH2-I) may also be obtained. These C-I dissociationenergies are in good agreement with other values for primary C-I bonds.2 The bonddissociation energies are given in table 3.TABLE 3 .-BOND DISSOCIATION ENERGIES DETERMINED IN THIS WORKbond D/kJ mol- bond DlkJ mol -Me,SiCH,-H 415+ 5 Me3SiCH2-I 227+ 6Me3CCH,-H 4 1 7 1 5 Me3CCH2--I 229+ 6NON-OBSERVED PROCESSESApart from intrinsic interest these studies provide some valuable informationabout several processes which were not observed.6(i) THE DISPLACEMENT REACTION, I=+Me4Si -+ Me,SiI+ Me-Had such a process occurred the reaction products would have been Me,SiI andmethyl iodide and complete consumption of iodine would have resulted.The obser-vation of a steady state and the failure to observe these products implies that thisprocess is unfavourable. Since the kinetics of reaction would be similar (apart fromreversibility) but no disturbance of the steady state occurs until temperatures > 680 Kwe can set a conservative limit of k6 < k , . In the middle of the observed rangeat T = 629 K, k , = 47.0 dm3 mol-l s-l and thus k6 < 0.47 dm3 mol-l s-l.If weassume A6 = dm3 mo1-l s-' by analogy with the A factor for the only othersimilar displacement process for which Arrhenius parameters have been measured,20uiz. ,7I-+ Me,SiSiMe, Me3SiI + Me,Si.,then E6 2 103 kJ mol-l. Since E7 = 33.9 kJ mol-', it is clear that it is considerablymore energetically difficult for an iodine atom to displace a methyl group from asilane than for it to displace a trimethylsilyl group. Part of this activation energ2914 GAS PHASE REACTION OF 12+Me4Sidifference is due to step (6) being more endothermic than step (3, the enthalpy dif-ference being exactly equal to the extra strength of the C-Si bond over the Si-sibond. The results of our earlier work on D(Me,Si-H) imply C-Si bond dissocia-tion energies of -380 kJ mol-1 while Davidson and Howard have obtained D(Me3-Si-SiMe,) = 337 kJ mol-l.Thus the enthalpy difference is only -40 kJ mol-lwhereas the activation energy difference is at least 70 kJ mol-l. This means that, inthe exothermic direction, displacement of iodine in trimethylsilyl iodide is morereadily achieved by trimethylsilyl radicals than by methyl radicals with an activationenerw difference of at least 30 kJ mol-1 favouring the former. This may well be thecase, (arising possibly because of some energy lowering attraction between incomingMe3Si* and the I atom in Me,SiI) but in many circumstances, such as hydrogenabstraction reactions, methyl radicals are undoubtedly more reactive than trimethyl-dyls.Thus it would be very helpful to have direct experimental corroboration ofthis point and indeed to uncover further examples of such displacement processes ofwhich very few are known on silicon.*8(ii) THE RADICAL REARRANGEMENT, Me,SicH, + Me2QiEtThis process, involving a methyl shift, has no parallel in hydrocarbon chemistry.It has not yet been observed and indeed several other potential rearrangements ofa-silyl alkyl radicals to silyl radicals have been found not to occur under the condi-tions ~tudied.~ 3-2 Nevertheless, such processes are not entirely implausible, owingto the availability of the silicon d-orbitals which might stabilise transition statesrelative to their hydrocarbon counterparts. The purpose of this discussion is simplyto use the experimental data obtained here to set a lower limit on its activation energy.Had such a process occurred, it would have led to ultimate formation of dimethyl-ethylsilyl iodide rather than the observed product.Again complete consumption ofiodine would have been expected. Failure to observe this, even at higher tempera-tures,' means that we can again set a limit of a factor of disfavouring its occur-rence in competition with other observed processes, such as,Me3SiCH2 + HI + Me,% +I*.The rate constant for this process is -108.* dm3 mol-1 s-l by analogy with similaralkyl radical abstractions from HI.2 At equilibrium, HI pressures are at least 1 Torrtmol dm-3) and so ks < x 108.8 dm3 mol-l s-l x 10-4-6 mol =102.2 s-l . If a reasonable A factor of 1012s5 s-l is assigned to this rearrangement(implying a small loss of entropy in the transition state due to restriction of motionof the migrating methyl group) then E8 124 kJ mol-l.Thus it is apparent thatdespite the availability of the silicon d-orbitals this reaction obviously has a highenergy barrier. The stability of this and similar radicals 23-25 is clearly kinetic inorigin and therefore nothing to do with p x - d, bonding favouring a-silyl radicalsthermodynamically as has been claimed. 149 2429(iii) THE '' GROUP SWITCHING " REACTION Me3SiCH21 --+ Me2SiIEtThe unimolecular process in which an electronegative atom or group migrates* Apart from those discussed here, the only other exampIes of which we are aware are l 2CF; + Me.& + Me3SiCF3 + CHgand 21, 22He + Si2Hs -+ SiH4 + SiH;.-f 1 Tom = 133.2Nm--2A .M. DONCASTER AND R . WALSH 2915from carbon to silicon while concurrently a less electronegative substituent movesfrom silicon to carbon has been invoked as an explanation for the rearrangement ofseveral a-carbon-functionalised organosilanes.26* 27 Recently for some chloromethyl-silanes, the rearrangement has been shown to be bimolecular.28 Whatever its naturesuch a process cannot be occurring in our system, since even were the product over-looked by us it would result again in complete consumption of iodine which was notobserved. As a unimolecular reaction, our observations on the stability of ourequilibrium lead to a limit of kg < 10-6-9 s-l. A reasonable A factor would be A g 2:1012.5 s-l for this process (once again a small entropy loss in the transition state dueto restriction of the molecule by the migrating groups) and from this Eg 234 kJmol-l.This shows that there is a high energy barrier to such a process as far asiodide and methyl switching are concerned. Although such processes may be morefavoured where more electronegative groups migrate,26p 27 we would not expect anyof them to be particularly easy on the basis of this result.APPENDIXTHE INTEGRATED RATE EQUATIONProvided [Me,Si] > [I,] integration of the rate eqn (A) givesw-1- ([I210/CI21e) lnf2 = ck$fi = (Y +Ye)/(Y-Ye).fi = (Y," +YY,)/(Y,2 - w e >wherec = NY; +u"e/YeCV,2 - ~ 3where y 2 , y,2 and y t are [I2It, [I2I0 and [IJe respectively, andb = [Me,Si], N [Me,Si],.The only unknown in eqn (B) is k, which can be determined by least squares fit of thedata.THE COLLISION CALCULATIONSThese were simply carried out using the standard hard sphere collision result,z = 02J8nkT/p, where the terms have their usual signifi~ance.~~ Collision dia-meters were as follows; I-, 0.41 nm (assumed identical to Xe 30), Me,C, 0.55 nm 30and Me,Si, 0.617 nm (obtained by increasing 0 for Me,C by twice the difference instandard bond lengths 31 between Si-C and C-C).The authors thank the S.R.C.for research funds in support of this work.A. M. Doncaster and R. Walsh, J.C.S. Faraday I, 1976, 72, 2901.D. M. Golden and S. W. Benson, Chem. Rev., 1969, 69, 125.R. Walsh and J.M. Wells, J.C.S. Faraday I, 1976, 72, 100.R. Walsh and J. M. Wells, J.C.S. Faraday I, 1976,72, 1212.E. A. V. Ebsworth, Organometallic Compounds of Group IV Elements ed. A. G. MacDiarmid(Marcel Dekker, New York, 1968), vol. 1, chap. 1.J. A. Kerr, A. Stephens and A. J. Young, Int. J. Chem. Kinetics, 1969, 1, 339. ' A. U. Chaudry and B. G. Gowenlock, J. Organometallic Chem., 1969, 16,221.E. R. Morris and J. C. J. Thynne, J. Phys. Chem., 1969, 73, 2394; J. Organometallic Chem.,1969, 17, 3.R. E. Berkley, I. Safarik, H. E. Gunning and 0. P. Strausz, J. Phys. Chem., 1973, 77, 1734.l o T. N. Bell and A. E. Platt, J. Phys. Chem., 1971, 75, 603 ; Int. J. Chem. Kinetics, 1971, 3,307291 6 GAS PHASE REACTION OF I,+Me,Sil 1 W. J. Cheng and M. Szwarc, J.Phys. Chem., 1968,72,494.l 2 T. N. Bell and A. E. Platt, Int. J. Chem. Kinetics, 1970,2,299.l4 P. J. Krusic and J. K. Kochi, J. Amer. Chem. Soc., 1969, 91, 6161.l6 J. A. Kerr, Chem. Rev., 1966, 66,465.l7J.A.N.A.F. Thermochemicul Tables, ed. D. R. Stull and H. Prophet, NSRDS-NBS 37, (Nat.l8 S. W. Benson, Thermochemicul Kinetics (Wiley, New York, 1968), pp. 30-42.l9 A. M. Doncaster and R. Walsh, to be published.2o S. J. Band and I. M. T. Davidson, Truns. Furuduy Sac., 1970, 66,406.21 K. Obi, H. S. Sandhu, H. E. Gunning and 0. P. Strausz, J. Phys. Chern., 1972, 76,3911.22 T. L. Pollock, H. S. Sandhu, A. Jodhan and 0. P. Strausz, J. Amer. Chem. SOC., 1973,95,1017.23 I . M. T. Davidson and C. A. Lambert, J. Chem. SOC. (A), 1971, 882.24 J. W. Wilt, 0. Kolewe and J. F. Kraemer, J. Amer. Chem. Suc., 1969, 91, 2624.2 5 K. Yamamoto, K. Nakamishi and M. Kumada, J. Orgunometullic Chern., 1967,7,197.26 W. I. Bevan, R. N. Hazeldine, J. Middleton and A. E. Tipping, J.C.S. Dalton, 1975, 252, 620.27 A. R. Bassindale, A. G. Brook, P. F. Jones and J. M. Lennon, Cunud. J. Chern., 1975, 53,332.28 J. M. Bellama and J. A. Morrison, J.C.S. Chem. Comm., 1975, 985.29 See, for instance, G. L. Pratt, Gas Kinetics (Wiley, London, 1969), p. 88.30 S. C. Chan, B. S. Rabinovitch, J. T. Bryant, L. D. Spicer, T. Fujimoto, Y. N. Lin and S. P.31 B. Beagley, Chem. SOC. Spec. Period. Rep. Moleculur Structure by Diflraction Methods, 1973,A. F. Trotman-Dickenson, Chem. and Ind., 1965, 379.H. Sakurai, A. Hosomi and M. Kumada, Bull. Chem. SOC. Japan, 1971, 44, 568.Bur. Stand., Washington, 2nd edn., 1971).Pavlou, J. Phys. Chenz., 1970, 74,3160.1,59, 111.(PAPER 6/788

ISSN:0300-9599    

DOI:10.1039/F19767202908

出版商:RSC    

年代:1976

数据来源: RSC

摘要:

Testing Intermolecular Potential Functions usingTransport Property DataPart 3.-Binary Diffusion Coefficient of Methane + PerfluoromethaneBY ANTHONY A. CLIFFORD AND ERIC DICK INS ON*^Department of Physical Chemistry, University of Leeds, Leeds LS2 9JTANDG. PETER MATTHEWS AND E. BRIAN SMITHPhysical Chemistry Laboratory, South Parks Road, Oxford OX1 342Received 26th April, 1976The binary diffusion coefficient DI2 for methane+ perfluoromethane has been measured by atwo-bulb method at 302.8, 323.2 and 342.9 K, and the application of a simple I' inversion " techniquegives three points on an effective spherically averaged intermolecular pair potential. Comparisonwith the spherical CHs-CF4 potential obtained from a full inversion of interaction viscosities indi-cates differences in the separation at a given energy of the order of 2 %.Of the three main transport coefficients of a binary mixture-viscosity, diffusionand thermal conductivity-only the mutual diffusion coefficient D1 is determinedprimarily by the cross term interactions.On this basis, therefore, D12 should be oneof the most useful non-equilibrium sources of information about intermolecularforces between unlike species. Unfortunately, however, it is difficult to measureaccurately and has been determined for most systems with a precision of only 1-2 %at the very best1In this work, DI2 for methane+perfluoromethane is reported at three tempera-tures using N,+H, as a standard system to calibrate the apparatus.2 CH4+CF4is the simplest example of a mixture of an aliphatic hydrocarbon with an aliphaticperfluorocarbon, a type of binary system that has been the subject of many thcrmo-dynamic investigations due to an interest in the cross interaction energy, which isanomalously weaker than that predicted from the normal combining rules.Sincethe viscosities qmix of CH,+CF4 mixtures have recently been measured over a widetemperature range,4* we are now in a position to compare the extent to which theDI2 and qmiX data can be represented by the same spherically averaged intermolecularpair potential.DETERMINATION OF Dl2The two-bulb apparatus is based upon the design of van Heijningen et al; itwill be described in detail elsewhere.'In brief, a composition difference Ax is set up at constant temperature betweentwo chambers connected by a capillary tube of known length and diameter, and arelaxation time is determined by monitoring the changing difference between thet Present address : Physical Chemistry Laboratory, South Parks Road, Oxford OX1 3QZ.291 291 8 CH4-CF4 POTENTIAL FUNCTIONthermal conductivities of the gases in the two chambers.The conductivity changesare followed in situ by two self heating thermistors, one in each chamber, which formpart of a sensitive bridge circuit.For a given pressure p and mean composition x, Oi2 is determined from a plotof ln(Ax(t)-Ax(t+$T)} against the time t , where T is the total time over whichmeasurements are taken. This procedure avoids the difficulty of needing to knowAx@) or ( h x ( t ) - A x ( ~ ~ ) } , each of which is difficult to measure because of long termelectrical drift and the necessity for establishing a quasi-stationary state.Correctionsare applied for the finite volume of the connecting tube and the difference betweenits effective and geometrical length^.^r3'(a) 0 . 8 2 2 0.802 4 6 8p-llatrn-'FIG. 1.-The pressure dependence of D12 for N2+Hz. D12p is plotted against l/p (tube length30.00mm, i.d. 2.06mm): (a) 302.8K; (b) 323.2K; (c) 342.9K.Much of the total error in the two-bulb experiment is associated with uncertaintiesin the geometry of the apparatus, in particular the diameter of the connecting tube.It is useful therefore to calibrate the apparatus with a standard system for which veryaccurate data have already been obtained. We choose N2 + H2 since Carson et aZB2have recently measured Di2 for this system near room temperature to a stated pre-cision of 0.2 %.Fig. 1 shows plots of our values of Di2p against l/p for the equi-molar N2 +H2 mixture at 302.8, 323.2 and 342.9 K (tube length 30.00 mm, i.d.2.06* mm). Within the experimental scatter, D I 2 p is independent of pressure at eachtcrnperature, thus indicating the absence of Knudsen effects.6TABLE 1 .-BINARY DIFFUSION COEFFICIENT DI2 OF CH4+ CF4D12(1 atm, x = 0.5)/cm2 s-1literature-this work aT/ K302.8 0.148 0.126"0.145120.1435323.2 0.167342.9 0.186Based on Ol2 for N2+Hz at 300 K estimated from ref. (2) ; the uncalibrated values are higherby 0.7 %. b Extrapolated to 302.8 K.C Calculated from mixture viscositiesA . A . CLIFFORD, E . DICKINSON, G . P . MATTHEWS AND E. B . SMITH 2919In fig. 2, uncorrected values of D12p for CH,+CF, (x = 0.5) at 302.8, 323.2 and342.9K are plotted against I/'. DI2p is independent of pressure in the range0.13-0.35 atm, but above -0.4 atm the data become spurious, probably dueto the presence of convection.1o The increased scatter at 342.9 K is caused by0.1851 Ip-I latm-'FIG. 2.-The pressure dependence of D12 for CH4+CF4. D12p is plotted against 1/p (tube length30.00 mm, i.d. 2.06 mm) : (a) 302.8 ; (b) 323.2 K ; (c) 342.9 K.a lower thermistor sensitivity and poorer temperature control. The values of D12(1 atm, x = 0.5) listed in table 1 are based upon the calibration system N2 +H2, forwhich Ol2 (1 atm, x = 0.5) = 0.803, cm2 s-l at 300 K.The result at 302.8 K iscompared with previous literature values. 99INVERSION OF VISCOSITIESViscosities of gaseous CH,+CF, mixtures have been reported as a function ofcomposition by Maitland and Smith (295-1022 K) and Gough et aL5 (150-320 K).The two sets of experimental ymix data agree well in the overlap region near 300 K,but the interaction viscosities y12 derived from them are inconsistent due to differentvalues being taken for the viscosity of pure CF4. Because of high slip-flow effects,Maitland's value is the more likely to be in error, so in order to remove artefacts inthe final potential which would arise from such a discontinuity in y12, the Maitlandvalue for pure CF, is reduced by 1.2 % to coincide with that of G o ~ g h .~ Values ofthe parameterATz =where Sz!'; ')* and 2)* are respectively the cross term collision integrals appro-priate to diffusion and viscosity, are taken from the Barker-Bobetic-Maitland-Smith(BBMS) potentia1,l and these are then used to recalculate interaction viscositiesfrom the adjusted qmix data. (The function chosen here is not crucial, since a 3 %change in AT2, the maximum range observed for all the potentials considered in thiswork, alters q12 by no more than 0.5 %. In fact, values of AT2 were generated fromthe final viscosity potential listed in table 2. The difference in q12 is 0.2 % at the lowesttemperature rising to 0.5 % at the highest, well within the experimental uncertainty of+2 %.2920 CH4-CF4 POTENTIAL FUNCTIONThe interaction viscosities are " inverted " as described previously l4 for purecomponent viscosities.In essence, the method seeks to relate each value of thecollision integral obtained from the experimental data to the coordinate of a pointon the potential energy function at an energy defined, initially, by an approximatepotential energy function. The derived function can be iteratively refined if reason-TABLE 2.-CH4-CF4 POTENTIAL OBTAINED FROM INVERSION OF INTERACTION VISCOSITIESr/nm0.35410.35810.36210.36610.37000.37400.37800.38200.38600.38990.39380.39570.39880.40040.40240.4047{UWkIlK31 30.02547.02052.01633.01280.0984.0736.2529.9359.2218.7114.247.3- 10.3-31.8- 52.5- 72.8rinm0.40740.41060.41430.41860.42350.42910.43520.43680.43840.44010.441 70.44330.44480.44630.4477{U(r)lkl/K- 92.4-111.5- 130.0- 147.7- 164.3- 179.4- 191.9- 194.4- 196.6- 198.3- 199.5- 200.0- 199.4- 197.6- 195.7r/nm {UW/kllK0.4489 -193.30.4508 - 190.50.4533 - 187.20.4561 -0.4591 -0.4623 -0.4657 -0.4695 -0.4736 -0.4780 -0.4827 -0.4879 -0.4933 -0.4992 -0.5054 -83.579.374.669.664.058.151.745.037.830.122.013.4rlnm0.51 190.51860.53850.55840.57830.59820.61810.63800.65790.67780.69770.7 1760.73750.99521.1942{UWlk 1 /K- 105.4- 98.0- 79.0- 63.9- 52.0- 42.5- 35.0- 29.0- 24.1- 20.2- 16.9- 14.3- 12.2- 2.0- 0.7r /nmFIG.3.-The CH4-CF4 intermolecular potential energy U(r) as a function of the separation r : -,two-iteration inversion of q12 ; 0, simple inversion of DI2.able assumptions are made about its behaviour beyond the range defined by experi-ment. Preliminarily, several long and short range extrapolations were tried, and thenumber of iterations was varied. The data were inverted assuming a range of valuesof the well depth E , and in each case the fit to the interaction second virial coefficientA. A . CLIFFORD, E . DICKINSON, G . P . MATTHEWS AND E . B. SMITH 2921of Douslin et aZ.15 was examined. Depending on the choice of extrapolation andnumber of iterations, the optimum value of ~ / k lay somewhere in the range 180-220 K :this represents the limit of precision that could be obtained from the macroscopicinformation alone.To obtain a more precise estimate of E it is necessary to make an assumption aboutthe behaviour of the intermolecular energy function at long range.There are notheoretical estimates of the r6 or r-8 terms in the multipole expansion for CH4-CF4,and so we have assumed that, since the CH4-CH4 interaction is closely representedby a (20 : 6) function,16 a moderately conformal CH4-CF4 potential will also be reason-ably represented by a suitably scaled (20: 6) function in the long range attractiveregion. The final potential is insensitive to the short range extrapolation, but a one-term Barker function l 3 is chosen as this was found to be the most satisfactory forThe final potential is listed in table 2 and illustrated in fig.3. It is obtained aftertwo iterations from a BBMS starting potential as further iterations produced noimprovement in the fit to the viscosities. The form is independent of the startingpotential within -1 % in r, the viscosities and virial coefficients being fitted withr.rn.s. deviations of 1.1 x lo-' kg m-l s-I and 1.9 cm3 mol-1 respectively. Thecharacteristic parameters are ~ / k = 200 K and CT = 0.3975 nm, and the ratio of thewell depth to that of CH4-CH4 (0.922) is in good agreement with the ratio of the Boyletemperatures (0.91 7).15CH4-CH4.INVERSION OF DIFFUSION COEFFICIENTSSince the D I 2 values cover a relatively narrow temperature range, a full inversionanalogous to that of y12 is not feasible.We employ instead the simplified inversionprocedure described by Clancy et al. ;17 this is effectively a one-iteration inversionfrom the BBMS starting p0tentia1.l~ While it would be preferable to estimate Eindependently of the vl2 analysis, we can do no better than assume Elk = 200 K.The result of this inversion is shown in fig. 3. In terms of the separation r, the pointsobtained from inverting D12 differ from the viscosity potential by 1.8-2.2 %, with theformer having the smaller intermolecular separation.The difference between the two potentials is explicable in terms of experimentaluncertainties in D1 and q (- & 2 % each) and ignoring errors inherent in the analy-sis itself.However, these latter errors may be significant here since the inversion ofdata for CH4+ CF4 mixtures differs in two important respects from most of thosepreviously considered : (i) Each molecule contains 12 internal degrees of freedom,whilst the kinetic theory upon which the inversion method is based is properly validonly for monatomic species. (ii) The actual anisotropic pair interaction may perhapsbe represented only rather crudely by a potential of spherical symmetry. Addition-ally, there is the fact that the quantity y12 is not defined uniquely by the mixtureviscosities ; the pure component values and some approximate knowledge of thepotential form (through AT2) are also required.This work represents the first direct attempt to describe the form of the intermole-cular potential between hydrocarbon and perfluorocarbon molecules.Despite itsobviously approximate nature, we believe that this CH4-CF4 potential may usefullyaugment previous information obtained from the study of liquid mixtures. Thus,on the basis of deviations from the Berthelot combining rule suggested by liquid statethermodynamic measurement^,^ we can estimate, assuming conformality, the welldepth of the spherically averaged CF4-CF4 interactionE(CF~-CF~) = (E(CH~-CF~)/O.~O~)~/E(CH~-CH~) = (223 & 10) K2922 CH4-CF4 POTENTIAL FUNCTIONWe note, finally, that the refinement of potential energy functions obtained in thismanner, and the extension of the methods to other systems, is limited in the main bydeficiencies in the quality and range of the available thermophysical data.T.R. Marrero and E. A. Mason, J. Phys. Chem. Rex Data, 1972, 1, 1.P. J. Carson, P. J. Dunlop and T. N. Bell, J. Chem. Phys., 1972, 56, 531.J. S. Rowlinson, Liquids and Liquid Mixtures (Butterworth, London, 2nd edn., 1969).G. C. Maitland and E. B. Smith, J.C.S. Faraday I, 1974, 70,1191.D. W. Gough, G. P. Matthews and E. B. Smith, J.C.S. Faraday I, 1976,72,645.R. J. J. van Heijningen, A. Feberwee, A. van Oosten and J. J. M. Beenakker, Physica, 1966,32, 1649. ’ A. A. Clifford, E. Dickinson and R. S. Mason, to be published.* E. P. Ney and F. C. Armistead, Phys. Rev., 1947,71, 14.Lord Rayleigh, Theory of Sound (Macmillan, London, 1878), vol. 11, p. 295.lo A. A. Clifford, E. Dickinson and P. Gray, J.C.S. Faraday I, 1976,72, 1997.C. R. Mueller and R. W. Cahill, J. Chem. Phys., 1964, 40, 651.l 2 A. T. Hu and R. Kobayashi, J. Chem. Eng. Data, 1970, 15, 328.l3 G. C. Maitland and E. B. Smith, Mol. Phys., 1971, 22, 861 and references therein.l4 D. W. Gough, G. C. Maitland and E. B. Smith, MoI. Phys., 1973,25,1433.D. R. Douslin, R. H. Harrison and R. T. Moore, J. Phys. Chem., 1967, 71, 3477.l6 G. P. Matthews and E. B. Smith, to be published.l7 P. Clancy, D. W. Gough, G. C. Maitland, G. P. Matthews and E. B. Smith, Mol. Phys., 1975,30, 1397.(PAPER 6/801

ISSN:0300-9599    

DOI:10.1039/F19767202917

出版商:RSC    

年代:1976

数据来源: RSC

摘要:

Use of Labelled Propene to Distinguish between an Associative,a Dissociative and a Concerted Mechanism for theDouble Bond Shift Reaction of AlkenesBY CHRISTOPHER s. JOHN* AND CHRISTINE E. MARSDENDepartment of Chemistry, University of Edinburgh,West Mains Road, Edinburgh EH9 355RONALD DICKINSONANDChemistry Department, University of Glasgow, Glasgow G12 SQQReceived 26th May, 1976The specifically labelled propene, CD2=CH-CH3, has been used as a model alkene for investiga-tions of catalytic double bond shift reactions. Microwave spectroscopy has been used to analysedeuteropropene mixtures produced in this reaction. It is shown how in principle the positions ofthe labelling D atoms after reaction can be used to deduce whether an associative, a dissociative or aconcerted mechanism has operated.Specifically, the data for a-zeolite are consistent with an associa-tive mechanism, presumably on Brransted acid sites, and argummts are presented to show that thisis a general mechanism for alkene double bond shift reactions on this catalyst.Double bond shift reactions of alkenes may be envisaged as occurring by anassociative mechanism (gain of H+ by the alkene to generate a carbonium ion onthe catalyst surface), by a dissociative mechanism (loss of either H+ or H- fromthe alkene to generate a n-allylic species on the catalyst surface) or by a concertedmechanism (simultaneous H loss from the methyl group with H gain at the methyleniccarbon involving a cyclic type of surface intermediate). In the work to be describedwe use the simplest alkene, propene, specifically labelled as CD2=CH-CH3, as aprobe to deduce the operative mechanism.Conclusions are based on a comparisonof the product distributions which are predicted by these mechanisms as reactionproceeds with those experimentally observed. Each mechanism redistributes thelabelling D atoms around the molecules in a characteristic manner as is shown below.C H2= CH - C HD2 CD2 = CH-CHzD-DFIG. 1.-Dissociative mechanism for double bond shift of CD2==CH-CH3 ; -+ reaction path toprimary products, --+ reaction path to secondaryproducts. Electronic charges omitted for generality.2922924 MECHANISMS FOR DOUBLE BOND SHIFT OF ALKENESThe dissociative mechanism for double bond shift is shown in fig.1 where electroniccharges are omitted for generality; CD2=CH-CH3 is seen to dissociate a methylC-H bond, to produce the n-allylic species CD2-CH-CH, and, as the hydrocarbonis the only source of deuterium, this intermediate can initially only pick up H fromthe surface to regenerate the starting material or produce CHD2-CH=CH2, theprimary product. This [2H2]propene can lose D to the surface in a subsequentreaction, to produce CDH-CH-CH2, and gain H to give CDH2-CH=CH2 andCH3-CH=CHD in equal amounts, assuming random addition to either end of thesymmetrical n-ally1 species. Similarly, CD3-CH=CH2 and CDH2-CH=CD2should be formed in equal amounts by addition of D, produced in the previous reac-tion, to CH2-CH-CD,, the predominant n-ally1 species present on the surfaceinitially. Note that the [2H,’J- and [2H3]propenes are really secondary products, bothbeing produced by subsequent reaction of CHID2-CH=CH2._-------------____--_______I.____-_____________~-------..------------_______----.-.FIG. 2.-Associative mechanism for double bond shift of CD2=CH-CH3.The circled numbersshow the predicted concentrations of products produced from 10 molecules of reactant, based purelyon statistics, when hydrocarbon deuterium content is conserved.The associative mechanism is shown in fig. 2. A carbonium ion is produced byH+ addition from the surface. It can lose H+, to give CH2=CH-CHD2, in threeways whilst regenerating CD2=CH-CH3 in one way. It can lose D+ in two ways,to give only CDH=CH-CH3, considering only statistics and neglecting any kineticisotope effects.This provides D+ on the surface from the “start” of the reactionwhich can interact exclusively with the reactant to produce CD3-CH=CH2 as the~nlyl[~H,]propene present initially.s s5 -sFIG. 3.-Concerted mechanism for double bond shift of CD2==CH-CH3 ; --t reaction path to pri-mary products, -- --> reaction path to secondary products. S is a surface siteC . S . JOHN, C. E. MARSDEN AND R . DICKINSON 2925Provided the number of active sites on the surface is small (little dilution of pro-pene deuterium content as reaction proceeds so that deuterium balance requires equalamounts of [2H,]- and [2H3]propenes) not oiily are CH,=CH-CHD,, CH3-CH=CHD and CD3-CH=CH2 all primary products but the ratios of the amounts ofeach present are fixed as shown in fig.2.The concerted mechanism is shown in fig. 3. At the same time as a methyl C-Hbond is broken, and H lost to the surface site, H on an adjacent surface site is addedto the methylenic carbon atom and CD2H--CH=CH2 is produced as the primaryproduct. On subsequent reaction, CD2H-CH=CH2 regenerates the starting mater-ial and also produces CH3-CH=CHD as the only [2Hl]propene. Deuteratedsurface sites are produced and these will interact principally with the predominantdeuteropropenes present i.e. initially CD,=CH-CH3, to produce CD,-CH=CH,as the only [2H3]propene. The ['HI]- and [2H3]propenes produced by this mechanismare the same as those from the associative mechanism with the important differencethat they are secondary products, the ratios of amounts of CH2=CH-CHD,,CH3-CH-CHD and CD3-CH=CH, not being constant with time during theinitial course of the reaction.Different product selectivities are produced in the associative and concertedmechanisms compared with the dissociative mechanism.Arguments based onprimary and secondary products are used to deduce whether certain ratios will beconstant or will vary during the initial course of the reaction. These two factors allowall three postulated mechanisms to be distinguished in practice provided the reactionis clean, i.e., that deuterium exchange does not occur other than during isomerization.The isomerization of butene over 8-zeolite has recently been studied in theselaboratories as part of a broader programme to characterise this recently discovered,synthetic zeolite as its composition and structure suggest it should possess interestingcatalytic properties.2 Evidence has been acquired which indicates the involvementof carbonium ions in double bond shift and ring opening reactions.It was uncertain,however, whether these carbonium ions were produced by gain of H+ from the catalystsurface (associative mechanism) or by loss of H- to the surface (dissociative mechan-ism). It was desirable to use labelled propene to solve this problem, rather thanlabelled butene, as its chemical identity as a propene molecule is preserved followingdouble bond shift, only the position of the labelling being altered and easily detect-able using microwave spectroscopy as shown previously. This represents a consider-able experimental simplification over but- 1-ene, for example, which becomes adifferent species, a but-2-ene, following reaction which necessitates gas chroinato-graphic separation prior to the mass spectral analysis, used here to determine initiallythe extent of reaction.Perhaps more importantly, however, is the fact that all deutero-propene molecules are determined with equal sensitivity by microwave spectroscopy,to a first approximation, whereas deuterated cis-but-2-ene and but- 1 -ene have differingsensitivities due to different dipole moments and trans-but-2-ene, the major productover 8-zeolite,' is undetectable due to its zero dipole moment.EXPERIMENTALa-Zeolite (Shell Development Co.) was used in 130mg lots.Activation, known toinvolve both dehydration and loss of tetramethyl ammonium ions which were present duringsynthesis,lB 2 s consisted of heating in silica reaction vessels in uacuo (- 1 mPa) at 820 K for40 h, conditions known to optimise both catalytic activity and surface area (295 m2 g-l).Reactions were carried out at - 370 K, using 1.06 x lozo molecules of CD2=CH--CH3, andsamples were obtained for microwave analysis as described previ~usly.2926 MECHANISMS FOR DOUBLE BOND SHIFT OF ALKENESRESULTSGood agreement was found between the distributions of deuteropropenes calculatedfrom both mass and microwave spectra and D atoms were never found to be bondedto the central carbon atom in propene, both as found previ~usly.~ The deuteriumcontent of the hydrocarbon was found to remain constant within experimental errorduring reaction.TABLE 1 .-COMPOSITION OF [2H2]PROPENES AS A FUNCTION OF THE EXTENT OF REACTIONnormalised deuterium distribution/ %CD2= CH-CH3 CH2= CH-CHD2 CHD = CH-CH2D total [2H2lpropene/%82 18 065 34 155 45 010 30 60887865350 Equilibrium distribution with 5 exchangeable hydrogen atoms.TABLE 2.-cOMPOSITION OF [2H~]PROPENES AS A FUNCTION OF THE EXTENT OF REACTIONnormalised deuterium distribution/ %total [2H2]propene/ %CD€I=CH--CH3 CH2= CH-CH 2D87897510050401311250506088786510010035a Product distribution from operation of the associative or concerted mechanisms.b ProductC Equilibrium distribution with 5 ex- distribution from operation of the dissociative mechanism.changeable hydrogen atoms.TABLE 3.-cOMPOSITION OF C2H3]PROPENES AS A FUNCTION OF THE EXTENT OF REACTIONnormalised deuterium distribution/ %total[2H2]propene/CH2=CH--CD3 CDz=CH-CHzD CHD=CH-CHDz %91 9 0 8887 6 8 7873 18 9 65100 0 0 100 a50 50 0 10010 30 60 35 =0 Product distribution by operation of the associative or concerted mechanisms. b ProductC Equilibrium distribution with 5 exchange- distribution by operation of the dissociative mechanism.able hydrogen atoms.Table 1 gives microwave analyses for the composition of [2H,]propenes as a func-tion of the total quantity of [2H,]propene in the deuteropropene mixture.OnlyCHD2-CH=CH2 was produced and the data are most striking as a demonstrationof the " cleanliness " of the reaction ; no deuterium scrambling is evident and it isapparent that the great majority of molecules have undergone only a single surfacereaction. The data for [2H,]- and [2H,]propenes and the observation that no[2H,]propene was detected, even after 35 % loss of [2H2]propene, are also consistentwith a single surface reaction for most moleculesC, S.JOHN, C . E . MARSDEN AND R. DICKINSON 2927The data for [2Hl]propenes (table 2) indicate a high initial selectivity for CDH=CH-CH3 relative to CH2=CH-CH2D which is maintained as reaction proceeds.This provides good evidence for the operation of an associative or a concertedmechanism in double bond shift (fig.2 and 3 and table 2) and, conversely, evidenceagainst the operation of a dissociative mechanism. There is a notable differencebetween the [2H,]propene distribution after 35 % reaction of [2H2]propene fromthat predicted by statistics at equilibrium.The data for [2H3]propene (table 3) provide the same information as from [2H,]-propene but are even more convincing as a demonstration of the associative or con-certed mechanism, the initial selectivity to CD3-CH=CH2 being >90 %.DISCUSSIONThe product selectivities for both [,HI]- and [2H3]propenes (tables 2 and 3) showclearly that an associative or concerted mechanism is active in double bond shift ofpropene as (i), the agreement with predictions based on these mechanisms is strikingand (ii), the differences between observations and predictions based on the dissociativemechanism or on statistical equilibrium are considerable.It is informative to compute the ratios of the amounts of CD,H-CH=CH,,CHD=CH-CH3 and CD3-CH=CH2 produced as the reaction proceeded.Thesewere found to be 2.5 : 0.64 : 1.0 (12 % reaction), 2.7 : 0.80 : 1.0 (22 % reaction) and2.4 : 1.0 : 1.0 (35 % reaction) which are constant within experimental error, especiallyconsidering the very appreciable extent of reaction investigated, in agreement withprediction from the associative mechanism shown in fig. 2 and in complete disagree-ment with prediction from the concerted reaction. In addition, it is believed that theaverage value for the ratio of CD,H-CH=CH, to CDH=CH-CH3, of -3 : 1,differs from the predicted ratio, of 1.5 : 1, due to the operation of a kinetic isotopeeffect in the rate determining step which is therefore associated with rupture of aC-H or C-D bond in the carbonium ion.The ease of breaking a C-H bond com-pared with a C-D bond is deduced to be -3/1.5 or -2, a value which is reasonableboth with regard to predicted values from the known difference in zero point energiesof the bonds, values of <5, and to previous values deduced for reaction over MgOand Ti02.3It is concluded, therefore, that an associative mechanism is operative, that theactive site for this reaction is of the Brarnsted acid type and further that these sites arerelatively few in number because of the absence of an experimentally detectabledecrease in the deuterium content of the hydrocarbon.This result seems reasonablewhen the acid nature of this aluminosilicate, with a high silica to alumina molar ratioof ~ 8 , ~ is realised, and might therefore be expected to be a general result for mostacidic catalysts. Indeed, Weeks et a1.,2 working with a partially ammonium ionexchanged R-zeolite, reported the appearance of a fairly intense zeolitic hydroxy bandin the i.r. region at -3600 cm-l following activation at 773-873 K.It was shown previously that both magnesia and rutile catalysed double bond shiftreactions via the dissociative mechanism, presumably involving n-allylic carbanions,as the predictions of this mechanism (fig. 1) were observed; initial production ofCHD2-CH=CH2 from CD,=CH-CH, which was followed by the appearance ofthe predicted [2H,]- and [2H3]propenes.Here, over R-zeolite, clear evidence has beenobtained for the associative mechanism (fig. 2) ; simultaneous production of CHD2-CH=CH2, CHD=CH-CH3 and CD3-CH=CH2. We have clearly shown,therefore, the ability of CD2=CH--CH3 to distinguish between these mechanisms,a distinction that is very difficult to make in other ways. Investigations are underwa2928 MECHANISMS FOR DOUBLE BOND SHIFT OF ALKENESat present to decide whether a concerted mechanism is operative over an only slightlydehydroxylated y-A1203, as suggested by Knozinger et al.It seems reasonable that the mechanistic conclusion reached here with propeneover Q-zeolite is a more general one that can at least be carried over to butenes.At370 K it was found that the reversible first order rate constants for but-1-ene isomeri-zation (kb) and for [2H,]propene loss (kp) were related as kb = 8kp when compatibleunits were used. Fig. 2 shows on purely statistical grounds that if 10 molecules ofCD2=CH-CH3 react, then 6 of them will remain as [2H2]propenes whereas 4 willbe transformed to [2Hl]- and [2H3]propenes. Thus the statistical probability forloss of [2H,]propene in a single surface reaction is 2/5. Fig. 4 shows the situationfor but-1-ene isomerizing over Brarnsted acid sites. On statistical grounds, the prob-ability of but-1-ene loss in a single surface reaction can be predicted as 2/5 and sok, should equal kp, if the assumption that (a) both processes proceed through an iden-tical rate determining step and (b) the heats of adsorption are similar, are both valid.FIG. 4.-Associative mechanism for double bond shift of but-1-ene.The circled numbers show thepredicted concentration of products produced from 5 moleculzs of but-l-ene on a purely statisticalbasis.However, the operation of a kinetic isotope effect with propene of -2 suggests thatC-H bond rupture is involved in this rate determining step. As loss of [2H2]propeneis limited by the rate of breaking a C-D bond, whereas isomerization of but-1-eneis limited by the rate of breaking a C--H bond, we might now predict kb = 2k,. Inaddition it seems most reasonable that loss of H+ from the s-butyl carbonium ion willproceed more readily to produce the more stable but-2-enes, through secondaryC-H rupture, than to regenerate but-1-ene, through primary C-H rupture; afactor of 4 in rates from this source implies an activation energy difference for thebreaking of these bonds of 4.3 kJ mo1-l.The above two factors suggest stronglythat a ratio of k,/kp y 8 is consistent with identical mechanisms for the two reactions.Interestingly, the derived activation energies for the two reactions were Ep/kJ mol-Iof 35 1 and E,/kJ mol-l of 304 1, the difference again reflecting the kinetic isotopeeffect (an energy difference of -2.1 kJ mol-l) and the above activation energy differ-ence for primary with respect to secondary C-H bond rupture.It is of interest to consider the " cleanliness " of the overall reaction of CD2=CH-CH3 over Q-zeolite.As already commented, the data indicate that the majorityof molecules had undergone only a single surface reaction, even in the time requiredfor 35 % of [2H2]propene to have reacted. This is a consequence, amongst otherthings, of the associative mechanism. Consider that during a surface visit the prob-ability for desorption of an adsorbed molecule is much greater than that for surfacereaction; the probability of a molecule undergoing two surface reactions during asingle surface visit is therefore small. In the dissociative and concerted mechanisms,therefore, it would be necessary for CHD2-CH=CH2 to be formed during one sur-face visit and for it then to undergo a subsequent reaction before [2Hl]- and t2H3]C.S . JOHN, C . E . MARSDEN AND R . DICKINSON 2929propenes were produced ; we would necessarily require two surface visits before reac-tion was seen to proceed by mass spectrometry and so would expect to see the completespectrum of products produced by secondary reaction. Conversely, the associativereaction can proceed as seen by mass spectrometry following the very first surfacereaction of CD2=CH-CH3 (during which it is possible for D+ to be split off toproduce CHD=CH-CH3). Thus it is possible with an associative mechanism tohave 35 % of [2H,]propene reacted but with very few molecules having undergoneseveral surface reactions. In contrast, if the concerted mechanism had operatedthen after 35 % [2H,]propene loss, and with almost half that remaining being CHD2-CH=CH2 (table l), this species, in addition to CD2=CH-CH3 would have under-gone reaction with deuterated surface sites.This would have produced CH,D-CH=CDH and CH2D-CH=CD2 in almost equal quantities (assuming a kinetic isotopeeffect of 2). No CH2D-CH-CDH was detected in these experiments which isfurther proof for the associative mechanism being dominant over a-zeolite.In conclusion, we have shown that the specifically labelled propene, CD2=CH-CH3, may be used with advantage as a model alkene for investigation of double bondshift reactions with the redistribution of the labelling atoms, followed by microwavespectroscopy, being studied as the reaction proceeds. We have shown that differentproduct distributions can be used to differentiate between possible associative, dis-sociative or concerted reaction mechanisms and have deduced for a-zeolite that anassociative mechanism, occurring presumably on Brsnsted acid sites, is a generalmechanism for double bond shift in alkenes.(C. E. M.) thanks the University of Edinburgh for the award of a Vans DunlopScholarship and (R. D.) thanks the S.R.C. for a studentship. We all wish tothank Prof. C. Kemball, Dr. H. F. Leach and Dr. J. K. Tyler for their interestin this work and in particular we are indebted to Dr. J. K. Tyler for the provisionand use of a microwave spectrometer.The gift of the a-zeolite from Dr. J. F. Cole (Shell Development Co.) is gratefullyacknowledged.H. F. Leach and C. E. Marsden, to be published.T. J. Weeks, Jr., D. G. Kimak, R. L. Bujolski and A. P. Bolton, J.C.S. Farada-y I, 1976, 72,575.C. S. John, C. Kemball, R. Dickinson and J. K. Tyler, J.C.S. Faraday I, 1976, 72, 1782.J. F. Cole and H. W. Kouwenhoven, Adv. Chem. Ser. 121, ed. W. M. Meier and J. B. Uytter-hoeven ( h e r . Chem. SOC., 1973), p. 583.H. Knozinger, A. Corado, Gy. Gati, H. Hiestetter, A. Kiss, R. Letterer and H. D. Miiller,Symp. on the Mechanisms of Hydrocarbon Reactions (Sibfok, Hungary, 1973), p. 333.(PAPER 6/1003

ISSN:0300-9599    

DOI:10.1039/F19767202923

出版商:RSC    

年代:1976

数据来源: RSC

摘要:

Ionization Constants of Phenols in Methanol+ Water MixturesBY COLIN H. ROCHESTER* AND DAVID N. WILSONChemistry Department, The University, Nottingham NG7 2RDReceived 27th May, 1976The pKa values of six 2-substituted phenols, three 3-substituted phenols and their 3,5-disubstitutedanalogues, 1-naphthol and 2,6-di-t-butyl-4-nitrophenol in methanol + water mixtures at 298 K havebeen determined either spectrophotometrically (0.05 < XMeOH < 0.6) or potentiometrically (0.7 <XMeOH < 0.9). The acidities of the phenols in methanol were also determined where literature valueswere not available. Values of the Hammett p-parameter have been deduced. The variation of p withsolvent composition differs for 2- and 3-substituted phenols. 2-Substituents caused steric interferencewith the solvation of phenoxide anions and thereby led to an acid weakening effect which became greateras the methanol content of the solvent was increased.The Hammett equationW a , s = - PO+P&,IZwhere pKa,s and0 independent of solvent for 3-substituted phenols in all methanol+ water mixtures.2-substituted phenols the equation was not strictly valid.refer to a substituted phenol and phenol itself respectively was applicable withHowever forIonization constants have previously been reported for phenol and five 4-substi-tuted phenols in methanol+water mixtures over the complete range of solvent com-position.' The results were consistent with the applicability of the Hammett porelationship for all solvent mixtures with a maximum value of p at -0.45-0.55 molefraction of methanol.The variations of pK, and p values with solvent could not beexplained solely in terms of electrostatic effects. The total medium effects wereapparently due to a combination of electrostatic and non-electrostatic contributions.The influence of solvent structure and specific solvation effects on the results werebriefly discussed. The present work has been concerned with testing whether theconclusions drawn from the study of 4-substituted phenols were equally applicablein relation to the ionization of 2-substituted and 3-substituted phenols in methanol +water mixtures.EXPERIMENTALMethanol and water were purified as before.2 Commercial samples of 2-nitro-, 3-nitro-,4-nitro, 2-methyl-, 2-methoxy-, 3,5-dichloro- and 3,5-dimethylphenol and 1-naphthol wererecrystallized to constant melting point.2-Chloro-, 2-fluoro-, 2-formyl-, 3-chloro- and3-methylphenol were purified by distillation under reduced pressure. Previously preparedsamples of 2,6-di-t-butyl-4-nitrophenol and 3,5-dinitro-4-aminophenol were recrystal-lized. 3,5-Dinitrophenol was prepared from 3,5-dinitroani~ole.~ Components of buffersolutions and standardized solutions of sodium hydroxide in water, sodium methoxide inmethanol, and hydrogen chloride in methanol were purified or prepared as bef0re.l. 2 s 6* 'The basic procedures which were used for making spectrophotometric and potentiometricmeasurements and for the subsequent calculation of pKa values have been described in detaile1sewhere.l. For mole fractions of methanol xM < 0.6 the extent of ionization of thephenols in KH2PO4/Na2HPO4/KCl, borax/KCl, or NaHC03/Na2C03/KCl buffer soh-tions * was determined with a Unicam SP8000 spectrophotometer.The pH* values ofbuffer solutions in solvent mixtures other than those studied by Bates et aL8 were deduced293C . H . ROCHESTER AND D. N. WILSON 293 1as before,' the procedure adopted for NaHC0,/Na2C03/KC1 buffer solutions being identi-cal to that followed for aqueous KH2PO4/NazHPO,/KC1 mixtures. For mole fractions ofmethanol xM = 0.7, 0.8 and 0.9 the pH* values of solutions containing mixtures of a phenoland sodium hydroxide were measured using a combined glass and silver/silver chlorideelectrode system in conjunction with a pH meter.l The pKa values of phenols in puremethanol were deduced from spectrophotometric measurements of the extent of ionizationof one phenol in buffer solutions containing methanolic sodium methoxide and either aceticacid or another phen01.~ 2-Nitrophenol and 3,5-dinitrophenol were ionized in buffer solu-tions containing acetic acid (pKa = 9.524 9* lo).Spectral changes for 2-nitrophenol andTABLE 1 .-pK, VALUES FOR FOURTEEN PHENOLS IN METHANOL+ WATER MIXTURES AT 298 Kmol fractionMeOH00.050.10.20.30.40.50.60.70.80.91 .o00.050.10.20.30.40.50.60.70.80.91 .o00.050.10.20.30.40.50.60.70.80.91 .osubstituent2-Me10.3310.3910.6411.0311.3111.5011.7211.9112.0912.3812.9114.902-CHO8.388.468.568.778.939.199.449.6810.0010.3410.9712.823,5-Me210.2010.3210.5410.8911.1811.4111.5611.7111.9212.2312.7814.622-OMe9.9910.1510.4110.781 1.061 1.2311.4211.5911.7812.1712.7914.482-NO27.237.287.447.657.827.958.288.498.789.179.8511.532,3-(cH)4Q9.399.539.7010.0610.3010.5710.8111.0111.2511.5512.0613.913-Me10.1010.2010.4010.7011.0411.2211.4411.5811.7212.0912.6614.483,5-C128.188.348.448.618.808.949.109.299.6310.0310.6112.112-F8.738.818.969.269.459.699.9410.1810.4010.7511.3312.943-C19.129.159.259.499.729.8710.0610.2410.6911.0611.6613.103,5-(NO2)26.736.796.836.967.167.157.437.497.738.088.6210.292-c18.56'5.648.779.059.299.559.7910.0410.3010.6811.2412.833-n028.368.518.608.788.959.109.269.489.8110.1910.7612.33 f2,6-(t-B~)2-4-N026.656.686.787.037.247.377.557.848.138.469.1310.89fa I-Naphthol ; b sources of data have been given elsewhere l3 ; C ref.(14) ; dref. (15) ; e ref. (1 6),see note in ref. (3) ; f ref. (11)2932 PHENOLS I N MeOH+H20for 1 -naphthol in buffer solutions containing 2-chlorophenol and for 2-nitrophenol in buffersolutions containing 2-fluorophenol enabled pK, values for 1 -naphthol, 2-chlorophenol and2-fluorophenol to be calculated. 3-Nitrophenol (pK, = 12.331 11) was ionized in buffersolutions containing 3-chlorophenol or 3,5-dichlorophenol, and 3.5-dinitro-4-aminophenol(pK, = 11.865 12) was ionized in solutions containing 2-formylphenol. 2-Methoxyphenolwas a sufficiently weak acid that the extent of its ionization in methanolic sodium methoxidesolutions could be measured dire~tly.~gAll ionization constants refer to a temperature of 25.0&0.1°C and to the mold activityscale with standard states in the particular solvent mixture for which a pK, value is quoted.RESULTS AND DISCUSSIONThe pK.values at 298 K for fourteen phenols in methanol + water mixtures overthe full range of solvent composition are listed in table 1. Of the fourteen phenolsonly the ionization of 2-nitrophenol 1 7 9 1 9 9 2o and 3-nitrophenol 17* 18* 2o in methanol+water mixtures appear to have been studied before.The agreement (fig. 1) betweenthe present results and the literature values confirms the reliability of the experimentalmethods adopted here.(b)I I I I I0-2 04 0.6 0.8 1Qmole fraction of methanolFIG. 1 .-p& values for (a) 3-nitrophenol and (b) 2-nitrophenol in methanol +water mixtures at298 K. 0 present work; 0 see table 1 ; V ref. (17); A ref. (18); A ref. (19); V ref. (20).THE BORN EQUATIONThe failure of the Born equation to give a satisfactory account of the acidity of4-substituted phenols in methanol + water mixtures is also apparent in relation to thC . H . ROCHESTER AND D . N. WILSON 2933ionization of 2-substituted phenols and 3-substituted phenols.Two examples aregiven in fig. 2. Born theory leads to the equationpK,(S)-pK,(M) = 121.6(~-o.0128)(G+n) 1 1relating the ionization constants Ka(S) and Ka(W) of a phenol in solvent S (dielectricconstant D) and water, respectively, to the radii rH and rA of the proton and thephenoxide anion respectively. The predicted linearity of plots of pKa(S) - pK,(W)against (1/D) was not observed for any of the phenols which have been studied.I 1 I I103f(l /D)-0.0128]dashed curve is for unsubstituted phenol.'FIG. 2.-Test of the Born equation for (a) 3,5-dimethylphenol and (b) 3,5-dinitrophenoI. TheFurthermore the results for all the phenols with electron withdrawing substituents,particularly for 3,5-dinitrophenol [fig.2(b)], were completely inconsistent with eqn (1)since no reasonable estimates of rH and rA predicted a solvent effect on pKa assmall as the experimental result. Thus, as before,l* * the medium effect for phenolswith strongly electron withdrawing substituents must contain an appreciable negativenon-electrostatic contribution to pKa(S) - pK,(W) in methanol + water mixtures.Eqn (1) predicts that increasing the radius r A of the phenoxide anion should decreasethe magnitude of the medium effect. Thus the medium effect for a substituted phenolshould always be less than that for phenol itself assuming that the substituent increase2934 PHENOLS IN MeOH+H,Othe effective radius rA of the acid anion. This expectation is not consistent with theresults for 3,5-dimethylphenol (fig.2), 3-methylphenol, 2-methoxyphenol and 2-methylphenol. The failure of the simple Born model in this context is apparentlyrelated to the electronic effects of the substituents in the phenol ring. Electronwithdrawing substituents give a smaller variation of pK,(S) with increasing methanolcontent of the solvent and electron donating substituents give a larger variation ofpKa(S) (compared with that for phenol itself) than is predicted by the theory. Eqn(1) leads to eqn (2), which compares the medium effects predicted by simple Borntheory for a substituted phenol (subscript s) and phenol itself (subscript u).[pK,(S)-pKa(W)],-[pK,(S)-pK,(W)], = 121.6 - -0.0128 -- - (A )(r:s *:I* (2)This equation is not applicable to the present results because for all the phenols studiedthe left hand side of eqn (2) passed through a maximum at XMeOH N 0.4-0.5.Thusthe non-linearity of the plots in fig. 2 cannot be solely attributed to the non-linearityas a function of (1/D) of the free energy of transfer of the proton from water tomethanol + water mixtures.22 Electrostatic theories, based on the Born model, whichsuggest that the medium effect on pKa should be a linear function of (l/D) 23* 24 orthat the relative medium effects of phenol and a substituted phenol should be linearin (1/D) 2 5 are also incompatible with the present results.THE HAMMETT EQUATIONThe applicability of the Hammett equationpK,(S), = - pa +constantto the ionization of 2-substituted phenols in any methanol+water mixture was con-firmed by the linearity of plots of pK, for the phenols in a given mixture against pKafor the same phenols in water.Hence p as a function of solvent composition (table 2)was deduced from the slopes of the lines and taking p = 2.23 for water solvent.26Values of p for 3-substituted phenols and 3,5-disubstituted phenols were deduced inthe same way (table 2). The two groups of phenols conformed to the same plotssince the increment to the pK, of phenol in any solvent mixture produced by a single3-substituent (Me, C1 or NO2) was doubled by the corresponding 3,5-disubstitution.TABLE 2.-vARIATION OF THE HAMMETT P-PARAMETERS, p2 FOR 2-SUBSTITUTED PHENOLS ANDp3 FOR 3-SUBSTITUTED PHENOLS IN METHANOL+ WATER MIXTURES AT 298 Kmole fractionMeOH P20 2.230.05 2.290.1 2.350.2 2.460.3 2.520.4 2.54mole fractionP 3 MeOH P2 n32.23 0.5 2.52 2.682.23 0.6 2.47 2.632.33 0.7 2.43 2.642.48 0.8 2.41 2.622.56 0.9 2.33 2.632.67 1 .o 2.42 2.77The variations of p with solvent for 2- and 3-substituted phenols are compared infig.3 with the results for 4-substituted phenols. The results for 2- and 4substitutedphenols are closely similar for all the solvents studied except for pure methanol wherep 2 > p4. However for all methanol+water mixtures p3 > p2 = p4 and thereforethe effect of solvent on the increment to pKa produced by a substituent was greatestwhen the substitution occurred in the 3-positionC . H . ROCHESTER AND D. N. WILSON 2935Comparison of the influence of solvent on the effects of 2- and 3-substitution onthe pK, of phenol shows that steric factors involving 2-substituents contribute to theresults.This is best exemplified by the plots in fig. 4 of pK, for the phenols inmethanol against pK, for the same phenols in water. The lines for the 2- and 3-substituted phenols not only differ in slope in accord with p2 # p3 but are also dis-placed in the sense that the decrease in Ka in passing from water to methanol solventI I I I I0.2 0.4 0.6 0.8 1 .omole fraction of methanolFIG. 3.-Variation of the Hammett p-parameters with solvent composition for (a) 2-substitutedphenols, (6) 3-substituted phenols and (c) Csubstituted phenols. Curves are staggered by 0.6 unitson the ordinate axis for clarity.is greater for a 2-substituted phenol than for a 3-substituted phenol with the samepKa in water.The point for unsubstituted phenol lies on the line for 3-substitutedphenols and not on the line for 2-substituted phenols. These effects are also apparentfrom the corresponding plots for the phenols in methanol + water mixtures althoughthe deviation between the two lines became less as the methanol content of the solventwas decreased. Methanol is a more bulky solvent than water and therefore the resultssupport the proposal that 2-substituents in a phenol impose steric effects on solvationprimarily of the phenoxide anions but probably also to a lesser extent of the neutralphenol molecules. Steric effects on solute-solvent interactions involving phenoxideanions have previously been identified by comparison of the acidities of methyl-phenols and dimethylphenols in water and methanol.16* ’’ The effects are en-hanced 3* 28 and have been characterized in some detail 2 9 9 30 for 2-t-butyl substitution.The present results suggest that 2-substituents other than alkyl groups give comparablesteric effects.In methanol + water mixtures steric effects become more significantwith increasing methanol content of the solvent.One corollary of the plots in fig. 4 is that if CT for a 2-substituent is taken as indepen-dent of solvent then the Hammett equationcannot be strictly valid for 2-substituted phenols in all methanol + water mixtures.The deviation of the point for phenol from the line for 2-substituted phenols impliesthat whereas eqn (3) is applicable, eqn (4) is not.The constant in eqn (3) cannot betaken equal to pKa(S),, except for the one unique solvent chosen as standard and forPKa(S)s = - PU + PKa(% (42936 PHENOLS IN MeOH+H,Owhich the a-values have been deduced. Tabulated a-values refer to water solvent.26Thus the deviation between eqn (3) and (4) for 2-substituted phenols is such that theconstant [eqn (3)] is greater than PK,(S)~ by an amount which increases as the methanolcontent of the solvent is increased. The deviations between the constant and pK,(S),reflect the solvent-sensitive steric effect of the 2-substituents on pK,. Eqn (4) was15 I I I14 -13 -s2 z 3 12-PKa(WFIG. 4.-pKa values in methanol as a function of the corresponding pKa values in water for 0 2-substituted phenols ; A 3- and 3,5-substituted phenols ; V phenol ; 4-substituted phenols.applicable both for 3- and for 4-substituted phenols.Thus for methanol + watermixtures plots of pK,(S) against pK,(W) for 4-substituted phenols were parallel tothose for 2-substituted phenols (p2 w p4) but shared the points for unsubstitutedphenol with the corresponding plots for 3-substituted phenols which were of steeperslope (p3 > p4). The only exception was for methanol solvent where the plots for2- and 4-substituted phenols were not parallel (fig. 4 ; p2 > p4).THE GRUNWALD Y- FUNCTIONThe present data confirm limitations which have already been exposed 31-33 inthe applicability of the Grunwald activity postulate 34 in relation to the ionizationof neutral acids in alcohol + water mixtures.Thus calculated values of Y-33- 34 forall the phenols studies were not a unique function of solvent composition. Valuesfor 2- and 4-substituted phenols differed at least in part because of the steric effects of2-substituents which also led to the non-applicability of eqn (4). Deviations betweenY- values deduced from the data for 4- and 3-substituted phenols are reflected in thedifferent trends in p for the two groups of acids as a function of solvent. Thus Y-should be a linear function of p with values of zero for water and 1.00 for purC . 13. ROCHESTER AND D. N. WILSON 2937methanol. It can be concluded from the present work that Y- may be a unique func-tion of solvent for a series of closely related acids such as 4-substituted phenols butthat this is not true when comparison is made of phenols with substituents in differentpositions in the benzene ring.EFFECTS OF 2,6-DI-t-BUTYL SUBSTITUTIONThe effects of solvent on the ionization of 2,6-di-t-butyl-4-nitrophenol were closelysimilar to the corresponding solvent effects on the ionization of 4-nitrophenol (fig.5).The difference between the pK, values of the two phenols in water was identical to thecorresponding difference for methanol solvent. The two curves in fig. 5 deviated byless than 0.15 pK, units over the entire range of methanol + water composition. Thisis a surprising result as the presence of two bulky t-butyl groups ortho to the phenolicgroup would be expected to influence solute-solvent interactions, particularly thoseinvolving the charged phenoxide anion.The fact that 4-nitrophenol is a weaker acidthan 2,6-di-t-butyl-4-nitrophenol is in itself unusual since the presence of t-butylgroups, particularly in the ortho position, generally decreases the acidity ofphenols.3* 6* 28 This effect has been discussed for methanol solvent for which it wassuggested that the thermodynamics of solvation of the 2,6-di-t-butyl-4-nitrophenoxidemole fraction of methanolFIG. 5.-Variation of pKa(S) - pKa(W) with solvent composition for 0 4-nitrophenol, A 2,6-di-t-but yl-4-nitrophenol.ion is dominated by solute-solvent interactions involving the nitro group.3* 29* 30Interactions between the charged phenolic oxygen atom and solvent molecules arehindered by the 2,6-di-t-butyl substituents.The charge on the anion therefore be-comes localized and stabilized on the 4-nitro group which can interact with solventmolecules without any steric hindrance from the t-butyl substituents. Differencesbetween the free energies of ionization of phenols are primarily influenced by thefree energies of solvation of the phenoxide anions rather than of the neutral phe2938 PHENOLS IN MeOH+H,On ~ l s . ~ ~ . 36 It is therefore reasonable that the medium effects of pK, for 2,6-di-t-butyl-4-nitrophenol and Ltnitrophenol should be similar if steric hindrance to solva-tion of the 2,6-di-t-butyl-4-nitrophenoxide ion is insignificant. It would be of interestto study a 2,6-di-t-butyl substituted phenol without a strongly electron withdrawing4-substitutuent but unfortunately such phenols are insoluble in methanol +watermixtures over the majority of the solvent composition range.THE ROLES OF SOLVATION AND SOLVENT STRUCTUREThe roles of solvation and solvent structure in the determination of the mediumeffects on pK, of 4-substituted phenols in methanol+water mixtures have been dis-cussed elsewhere.' The present results suggest that the arguments presented areequally applicable for 2- and 3-substituted phenols.The authors thank the S.R.C.for a studentship (D. N. W.).G. H. Parsons and C. H. Rochester, J.C.S. Furuduy I, 1975,71, 1058.G. H. Parsons and C. H. Rochester, J.C.S. Furuday I, 1972, 68, 523.C. H. Rochester and B.Rossall, J. Chem. SOC. (B), 1967,743.V. Gold and C. H. Rochester, J. Chem. SOC., 1964, 1727.N. V. Sidgewick and T. W. J. Taylor, J. Chem. SOC., 1922, 1853.C. H. Rochester, J. Chem. SOC., 1965, 676.C. H. Rochester, J. Chem. SOC. (B), 1967, 33.R. G. Bates, M. Paabo and R. A. Robinson, J. Phys. Chem., 1963, 67, 1833.T. Shedlovsky and R. L. Kay, J. Phys. Chem., 1956, 60,151.R. G. Barradas (Wiley, New York, 1966), p. 211.lo R. G. Bates and R. A. Robinson, in Chemical Physics of Ionic Solutions, ed. B. E. Conway andl1 P. D. Bolton, C. H. Rochester and B. Rossall, Truns. Furuduy Soc., 1970, 66, 1348.l 2 C. H. Rochester, Chem. undInd., 1971, 153.l3 C. H. Rochester, in The Chemistry of the Hydroxyl Group, ed. S . Patai (Interscience, London,l4 L. K. Creamer, A. Fischer, B. R. Mann, J. Packer, R. B. Richards and J. Vaughan, J. Org.l5 L. A. Cohen and W. M. Jones, J. Amer. Chem. SOC., 1963, 85, 3397.l6 C. H. Rochester, Truns. Furuday SOC., 1966, 62, 355.l7 R. Gaboriaud, Ann. Chim., 1967, 201.l9 R. A. Robinson and R. G. Bates, J. Res. Nat. Bur. Stand. A, 1966,70, 553. '* D. Jannakoudakis and J. Moumtzis, Chimika Chroniku, 1968, 33A, 7.21 E. E. Sager, R. A. Robinson and R. G. Bates, J. Res. Nut. Bur. Stand. A, 1964, 68, 305.22 J. Juillard, Bull. SOC. chim. Frunce, 1968, 1894.23 K. Hiromi, Bull. Chem. SOC. Japan, 1960, 33, 1251.24 R. Reynaud, Bull. SOC. chim. France, 1968,2279.25 H. Ohtaki, Bull. Chem. SOC. Japun, 1969, 42, 1573.26 G. B. Barlin and D. D. Perrin, Quart. Rev., 1966, 20, 75.27 C. L. de Ligny, Rec. Truv. chim., 1966, 85,1114.28 C. H. Rochester, J. Chem. SOC., 1965, 4603.29 C. H. Rochester and B. Rossall, Truns. Furaduy Soc., 1969, 65, 992.30 C. H. Rochester and B. Rossall, Truns. Furuday SOC., 1969, 65, 1004.31 R. Reynaud, Compt. rend. C, 1968, 267,989.32 R. Reynaud, Bull. SOC. chim. France, 1969, 699.33 R. Thuaire, J. Chim. ghys., 1970, 67, 1076.34 E. Grunwald and B. J. Berkowitz, J. Amer. Chem. Soc., 1951, 73,4939.35 G. H. Parsons, C. H. Rochester and C. E. C. Wood, J. Chem. SOC. (B), 1971,533.36 G. H. Parsons and C. H. Rochester, J.C.S. Perkin 11, 1974, 1313.1971), p. 327.Chem., 1961,26,3148.B. J. Steel, R. A. Robinson and R. G. Bates, J. Res. Nut. Bur. Stand. A, 1967, 71, 9.(PAPER 6/1014

ISSN:0300-9599    

DOI:10.1039/F19767202930

出版商:RSC    

年代:1976

数据来源: RSC