摘要:

J. Chem. SOC., Faraday Trans. I , 1982, 78, 1 165-1 176 The Decomposition of 2,2,3,3-Tetramethylbutane in KC1- and B,O,-coated Vessels in the Presence of Oxygen BY ROY R. BALDWIN,* MOHAMED W. M. HISHAM, ALAN KEEN AND RAYMOND W. WALKER Department of Chemistry, The University, Hull HU6 7RX Received 18th May, 1981 The decomposition of 2,2,3,3-tetramethylbutane (TMB) in the presence of oxygen has been studied in both KCI-coated and aged boric-acid-coated vessels. The values for k, obtained with the two types of vessel surface are in close agreement, and combination of the values over the range 400-542OC gives A, = 1.04 x 1017 s-l, El = 294.7 f 3 kJ mol-l, effectively identical with previous values using KC1-coated vessels, and thus confirming the thermochemistry suggested for the t-butyl radical : (CH,),C--C(CH,), --f 2(CH,),C.(1) The observed rate constant, kobs, defined by the equation -d [TMB]/dt = kobs[TMB] increases with increasing TMB concentration with both types of vessel surface and this variation has been used to evaluate the ratio k,/k$. The results with both types of vessel surface are in agreement, giving A,/& = 4.4 x lo5 (dm3 mol-' s-l)+, E4-8E7 = 81.7f 8 kJ mol-l over the range 400-520 OC: HO2 + (CH,),C-C(CH,), --* H20, +(CH,),C + C(CH,)zCHz (4) (7) HO, + HO, -+ H,O, + 0,. Previous papers' on the decomposition of 2,2,3,3-tetramethylbutane (TMB) in the presence of 0, have concentrated on the evaluation of the Arrhenius parameters for reaction (1) and on the use of the system to study the reactions of t-butyl radicals' with H,, D, and 0,, and to evaluate their entropy and enthalpy of formation., In KC1-coated vessels, where HO, radicals are efficiently destroyed at the vessel surface, the basic mechanism involves reactions (1)-(3), and gives the rate expression (i) - d [TMB]/dt = kl [TMB].(9 However, the observed rate constant increases2 both with increasing concentration of TMB and with addition of inert gas. This is caused by two contributions to a very short chain, from HO, radicals through reaction (4), which competes with reaction (3), and reaction (5), and from OH radicals through reactions (8), (9) and (5), the H,O, being formed by reactions (4) and (7) and being predominantly destroyed by the surface process (6). Reaction (10) also occurs to a very minor extent with k , , / k , = 0.01; Arrhenius parameters for this reaction have been given.(CH3),C-C(CH3), + 2t-C,H9 t-C,H, + 0, + i-C,H, + HO, surface HO, + $H,O+20, 11651166 D E C 0 M P 0 SIT I ON OF 2,2,3,3-TE T R A M E T H Y L BUT A NE (CH,),CC(CH,),CH, + i-C,H, + t-C,H, surface H,O, + H,O +p, HO, + HO, -P H,O, + 0, (7) The present paper uses the variation of the observed rate constant in eqn (i) with TMB concentration in both KC1-coated and aged boric-acid-coated vessels to evaluate the rate constant k,. EXPERIMENTAL Reactions were carried out in cylindrical Pyrex vessels, 20 cm in length and 5.1 cm diameter. KClcoatings were renewed every week, as previous studies2 have shown that the efficiency of the surface for destruction of HO, and of H,O, decreased markedly with older coatings. Aged boric-acid coatings were prepared as previously de~cribed.~ Studies were made over the range 400-520 "C in boric-acid-coated vessels but restricted to 420-470 OC in KC1-coated vessels.A wide range of mixture compositions (0.1-4.0 Torr TMB) was examined. The reaction was followed by using gas chromatography to measure the isobutene formed. Electromagnetic valves (opening and closing time < 0.1 s) were used with reaction times below 100 s to admit the gases from the mixing bulb into the reaction vessel, and to admit the reactants into the sampling bulb after a pre-determined time interval. To avoid complications resulting from reactions of isobutene, which become noticeable beyond 10% reaction, consumption of TMB was restricted to a maximum of 5% for those points used to elucidate rate constants.The time for 5% consumption varied from ca. 10 s at 520 O C to cu. 3000 s at 400 OC. RESULTS AND DISCUSSION KCl-COATED VESSELS If H,O, is always destroyed by reaction (6), and if the minor reaction (10) is ignored, the rate of reaction is given by eqn (ii) if reaction (3) is the sole termination process, and by eqn (iii) if reaction (7) is the sole termination process: - d [TMB]/dt = k, [TMB] + 2k1 k, [TMB]'/k3 (ii) - d [TMB]/dt = kl [TMB] + k4 (kl/k7)a[TMB]). (iii) If reaction (3) is the sole termination process, the chain length increases with increasing temperature, since E4 x 80 kJ mol-l, with increasing TMB concentration, and with inert gas addition (since k, is diffusion-controlled). If reaction (7) is the sole termina- tion process, the chain length increases with TMB concentration, is independent of inert gas and decreases as the temperature increases, since El (ca.290 kJ mol-l) > 2E4. Increase in temperature, in TMB concentration or addition of inert gas increases the HO, concentration and thus favours reaction (7) relative to reactions (3) and (4). Use of preliminary values, of k, and k,/k74 gives the percentage reaction due to the HO, chain for a mixture containing 2 Torr TMB+30 Torr 0, at pressures of 60 and 500 Torr (N, addition) shown in table 1 for a range of temperature. At 60 Torr, the percentage increases and then decreases at high. temperatures. This is because the second-order radical-radical reaction (7) increases in importance relative to theR.R. BALDWIN, M. w . M. HISHAM, A. KEEN AND R. w. WALKER 1167 first-order processes (3) and (4) as the temperature increases, due to the increased HO, concentration resulting from the increased rate of dissociation of TMB. Table 1 shows that the HO, chain contribution is small, and has its highest value in the range 400-470 OC. A further difficulty in using this small chain contribution to evaluate k, is that the OH chain becomes increasingly important as the temperature increases above ca. 450 OC; 2% of the total OH chain is due to reaction (lo), with k,,/k, 0.01, the remainder being due to the formation and decomposition of H,O,. TABLE VARIATION OF CHAIN LENGTH WITH PRESSURE AND TEMPERATURE chain reaction (%) HO, chain OH chain T/OC at 60 Torr at 500 Torr at 60 Torr at 500 Torr 400 7 31 2 2 440 12 29 2 6 470 16 24 3 16 500 12 12 5 31 540 9 5 14 60 To evaluate k , from the isobutene (IBE) against time curves, a computer treatment was used.It is convenient to express the HO, radical concentration in terms of the parameter G = k$[HO,], since it is the ratio R, = k,/k$ that is actually involved in the computer treatment. Similarly k, enters as the ratio R, = k,/k,k The differential equations for TMB, H,O,, O,, IBE, isobutene oxide (IBO) and H,O are given by eqn (iv)-(ix), and the stationary-state equations for HO,, OH and t-C,H, radicals are given by eqn (x)-(xii), respectively : - d [TMB]/dt = k, [TMB] + R, G[TMB] + k, [OH] [TMB] (iv) (v) d [H,O,]/dt = G2 + R4 G[TMB] - k6 [H20,] - k8 [MI [H202] - d [O,]/dt = (k, + klo) (t-C,H,] [O,] - 0.75R3 G - 0 3 , [H,O,] - G2 d [IBE]/dt = k, [t-C,H,] [O,] + R, G[TMB] + k, [OH] [TMB] d [IBO]/dt = klo[t-C4H,] [O,] d[H,O]/dt = k6 [H,0,] + 0.5R3G + k, [OH] [TMB] 2k, [TMB] + 2k, [MI [H,O,] = 2G2 + R, G 2ks [MI [H20214- kio [t-C,H,I Lo21 = kg [OH] [TMBI (vi) (vii) (viii) (ix) (x) (xi) 2k, [TMB] + R, G[TMB] + k , [OH] [TMB] = (k, + k1J [t-C,H,] [O,].(xii) G is obtained from eqn (x), and eqn (xi) and (xii) then give [t-C,H,] and [OH]. Eqn (iv)-(ix) are solved by the Kutta-Runge numerical integration method. H,O, rapidly reaches a quasi-stationary value, so that computer time can be considerably reduced by using a short time step until the H,O, is close to its quasi-stationary value, after which a stationary state is written for H,O, with a much bigger time step.Operation of the computer program requires values of the parameters k,, klo/k,, R,, R,, k, and k,. The yield of IBE at various times is calculated for a range of mixtures1168 DECOMPOSITION OF 2,2,3,3-TET R A ME THY LBU TAN E and compared with the experimental values. An optimisation procedure adjusts the parameters k , and R, so as to obtain minimum r.m.s. deviation. Reaction (10) plays only a very minor part, and the ratio k,,/k, = 0.01 obtained from previous studies, of the IBO yield is sufficiently accurate. k , is known from direct studies5', of the homogeneous decomposition of H,O,. The values of k , and k, depend on the efficiency of the surface for destruction of HO, and H,O,, respectively. In general, the surface destruction constant k, for a cylindrical vessel is given7 by eqn (xiii) k, = aA,/P( 1 + 4aB/P) (xiii) where A, = 32D,A/d2, B = 8D,A/~cd.D,A is the diffusion constant of the species in gas A at unit pressure, E is the surface efficiency for destruction of the species, c the average velocity of the species and d the vessel diameter. a is a coefficient expressing the variation of diffusion coefficient with mixture composition, and is given7? * by eqn (xiv) (xiv) where xA, xB and xc are the mole fractions of the components. The value of R, = k,/k,l has been obtained experimentally from a studys of the oxidation of HCHO in KCl-coated vessels, where the HO, concentration is given by eqn (x), but with 2k1, [HCHO] [O,] replacing 2k1 [TMB] : (1 1) HCHO + 0, + HCO + HO,. k,, was obtained by studies at very low HCHO concentrations, where the chain length is effectively zero.Since k, is the unknown parameters determining the rate of HCHO oxidation are R,, k, and k,,/k$ (12) A , [ = A , in eqn (xiii)] has been calculated9 for H,O, using the method given by Hirschfelder et aZ.;l0 use of 4N,) = 0.375 nm, a(H,O,) = 0.40 nm and a(0,) = 0.354 nm gives A , = 850 Torr s-' at 500 OC for N, in a 5.1 cm diameter vessel, and a ( 0 , ) = 0.989 relative to a(N,) = 1. Use of a(TMB) = 0.745 nm gives a(TMB) = 3.23. The value of A , varies with temperature, but the values of a are effectively constant. A value of kl,/k$ has been obtained,, by studying the effect of HCHO on the induction period of the H, + 0, reaction at 500 OC in 5.1 cm diameter aged boric-acid-coated vessels.The only unknowns are thus R, and B,, which relates k, to A , via eqn (xiii). Optimisation of the agreement between the observed and calculated value of the CO yield at various times for various mixtures at 500 OC gave R, = 4.56 (Torr3 s-l); and B, = 255 Torr (M = N,). The value of R, is very close to the value of 4.28 calculated from A , on the assumption that the diffusion constants for HO, are the same as those for H,O,. The difference of ca. 7% could be due to small errors in any of the quantities k,,lk,$ obtained from the H,+O, induction periods in the presence of HCHO, with an estimated error of f lo%, R, obtained from the HCHO+O, system with an estimated error of & lo%, the diffusion coefficient or the value (2.0 x lo9) taken fork,.Consequently, the experimental value of R, = 4.56 (Torr3 s-l)i at 500 OC has been taken, together with the calculated value of A, for H,O,. Values at other temperatures have been calculated from the change in diffusion coefficient with temperature given by the treatment of Hirschfelder et aZ.l0 The optimum value of B, [for H,O, in eqn (xiii)], and to a much lesser extent the value of R,, differ from those quoted earliers of 161 and 4.64, respectively. This is because the earlier results of HCHO+O, have been treated to allow for the interaction7 between the uniform profile of chain centres throughout the vessel, HO, + HCHO + H,O, + HCO.R. R. BALDWIN, M. w. M. HISHAM, A. KEEN AND R. w . WALKER 1169 resulting from homogeneous termination processes, and the diffusion profile created by an efficient surface termination.In the case of a homogeneous initiation process, 0, such as the decomposition of TMB, which is uniform throughout the vessel, this interaction has been studied' by solving the differential eqn (xv) by numerical integration methods : d2n dn Dy+(D/x)--6n2+0 dx dx = 0. n is the chain-centre concentration, x the vessel coordinate, 6 the mutual homogeneous destruction constant and 0 the initiation rate. Both with linear (first-order) homo- geneous termination, where algebraic solution is po~sible,~? l2 and with mutual (second- order) termination, the interaction may reduce the radical concentration by up to 20% below the value calculated assuming no interaction. Both in the HCHO+O, system and in the TMB+O, system, the interaction is further complicated by the fact that the dissociation of H,O, is a secondary source of initiation; since H,O, is formed, by reactions (4) and (7), from HO, which is not at a uniform concentration throughout the vessel, and since H,O, is predominantly destroyed at the surface [although by a process that is only moderately efficient at most pressures if B, = 255 Torr (M = N,)], the total initiation process is not uniform.This difficulty was overcome by solving the two stationary-state simultaneous differential eqn (xvi) and (xvii) d2n dn dt2 dx D,-+ (D,/x) -- Sn2 + 2kI [TMB] + 2k,[M]p = 0 d2P dP D,@+ (D,/x)z- k , [M]p+ 0.5 6n2 + k,n[TMB] = 0 (xvi) (xvii) where n = [HO,], p = [H,O,], and D,, D, are the corresponding diffusion coefficients.These equations were solved numerically by the method used to solve7 the single eqn (xv). With eqn (xv), it is necessary to guess no, the value of n at the centre of the vessel, to start the integration procedure, and the program varies this starting value until satisfactory convergence of successive solutions has been achieved. With a high ratio of homogeneous to surface termination, the initial value of no has to be located with increasing accuracy, otherwise unreal solutions are obtained. With eqn (xvi) and (xvii), initial values of both no and po have to be provided, and accurate estimates are required, particularly at high ratios of homogeneous termination to surface termination, so that several trials are often necessary. Since incorporation of this procedure into the main program, together with optimisation procedures, would make the running time excessively long, eqn (xvi) and (xvii) were solved separately, the volume average concentrations of HO, and H,O, were compared with the concentrations obtained on the assumption of no interaction, and corrections obtained which could be inserted into the main program.Because of the presence of mutual termination it was also necessary to obtain corrections to the volume average of [HO,],. The corrections to [HO,] varied from 0 to -2O%, to [HO,], from -25 to +25%, and to [H,O,] from - 15 to + 15%. However, although the accurate evaluation of the role of H,O, is complex, interpretation is not particularly sensitive to the parameters used for the surface destruction of H,O, since, under the conditions used to evaluate HO, + TMB, the contribution of the OH chain is reasonably small.Results obtained in this way are summarised in table 2. As table 1 shows, the contribution of the HO, chain is only some 20-30% even under the most favourable conditions. While the r.m.s. deviation between observed and calculated yields of IBE at 440 O C implies an accuracy of ca. 5% in k,/k,i, the1170 DECOMPOSITION OF 2,2,3,3-TETRAMETHYLBUTANE TABLE 2.-vALUES OF k4/k,i AT VARIOUS TEMPERATURES IN KCl-COATED VESSELSa r.m.s. no. of range of k4lk,$ dev. T/OC mixtures [TMB]/Torr /(dm3 mot1 s-l)$ (%) data ~ ~~~ 440 7 0.50-4.0 0.436 2.4 Evans2 41 8 8 0.25-4.0 0.257 8.0 Keen13 437.5 9 0.25-4.0 0.455 2.8 Keen13 466 10 0.25-4.0 0.674 6.2 Keen13 a In all cases the range of pressure was 60-500 Torr.results are less satisfactory at 418 and 466 O C . Moreover, Nalbandyanl* has reported the detection of HO, radicals when H,O, is decomposed on a KC1-coated surface. Such a possibility would increase the concentration of HO, radicals above the value calculated from the present mechanism, although it would be unimportant at low temperatures and pressures, since, very little H,O, is formed because HO, radicals predominantly undergo reaction (3). However, at higher temperatures and pressures reaction (7) becomes the dominant reaction, and formation of HO, radicals by surface decomposition of H,O, might become important. No indication appears available from Nalbandyan's work of the yield of HO, radicals produced in this way.BORIC-ACID-COATED VESSELS In aged boric-acid-coated vessels, the surface is extremely inefficients both for the destruction of HO, and of H,O,. To avoid the difficult treatment of the interaction of homogeneous and heterogeneous termination, as well as the possible complication of surface decomposition of H,O, to give HO, radicals, studies were made in aged boric-acid-coated vessels to see if results consistent with those in KC1-coated vessels could be obtained. Fig. 1 shows the [IBE] against time curves at 440 O C in a 5.1 cm diameter aged boric-acid-coated vessel for a mixture containing 1 Torr of TMB, 3 Torr of 0, and at various total pressures from 15 to 200 Torr, obtained by addition of N,. The curves are significantly autocatalytic and the extent of autocatalysis increases with pressure because of the M term in reaction (8).This contrasts with the effectively linear [IBE] against time curve over the first 5% of reaction in KC1-coated vessels, where the very low quasi-stationary concentration of H,O, is reached in a time that is usually small compared with the earliest sampling time, so that the initial autocatalysis is not detected. In aged boric-acid-coated vessels, the autocatalysis persists far into the reaction, well beyond the 5% consumption of TMB that was normally set as a limit to avoid secondary reactions of IBE. In boric-acid-coated vessels, as shown in table 3, the chain due to HO, radicals is significantly greater than in KC1-coated vessels at lower temperatures, although at higher temperatures the main termination process is reaction (7) with both types of surface.However, as table 3 also shows, the chain length due to OH is also greater because of the much higher yields of H,O,. Since the H,O, increased continuously with time in the early stages, values corresponding to ca. 2% reaction are given in table 3. Operation at low pressure, by reducing the value of M in reaction (8), reduces the OH chain. A pressure of 15 Torr of 0, + TMB was thus used, except at the lowestR. R. BALDWIN, M. w. M. HISHAM, A. KEEN AND R. w. WALKER 1171 time/min FIG. 1.-Effect of inert gas addition on [IBE] against time relationship. 1 Torr TMB, 3 Torr 0,, +N,, 440 OC. 0, N, = 11 Torr; x , N, = 26 Torr; A, N, = 56 Torr; 0, N, = 146 Torr; V, N, = 196 Torr. TABLE 3.-vARIATION OF CHAIN LENGTH WITH PRESSURE AND TEMPERATURE IN BORIC-ACID- COATED VESSELS For each condition, sample time selected corresponds to ca.2% reaction. TMB = 2 Torr. chain reaction (%) HO, chain OH chain at at at at at at T/OC 15Torr 60Torr 500Torr 15 Torr 60Torr 500Torr ~~ ~~~~ ~ ~ 400 45 37 25 18 35 61 440 35 32 18 8 19 62 470 21 19 11 8 18 61 500 15 14 10 6 13 511172 DE C OM POS I TI 0 N OF 2,2,3,3-TE TR A ME THY LB U T ANE temperature of 400 O C , when the partial pressure of 0, was 30 Torr and the total pressure was made up to 60 Torr by N, addition. An algebraic method of allowing for the OH chain can be devised since, in boric-acid-coated vessels, k, and k, are zero, and in the absence of an HO, chain, the combination of eqn (v) and (x) simplifies to eqn (xviii) -- fH2021 - k, [TMB].dt (xviii) For small consumption, this can be integrated to give eqn (xix) [H,O,] = k,[TMB] t . (xix) With the same assumption, and neglecting the 1 % contribution from reaction (lo), substitution of this expression for [H,O,] in eqn (xii) and eliminating [OH] between eqn (iv) and (xii) gives eqn (xx): - d [TMB]/dt = k , [TMB] + 2k8 [M]k, [TMB] t. (xx) 6.0 3.01 1 0 2 4 6 8 time/ min FIG. 2,-Variation of kobs with time in aged boric-acid-coated vessels. 440 O C , kobs = [IBE]/2[TMBlo t . 0, [TMB], = 4, [OJ = 11 TOIT; A, [TMB], = 2, [O,] = 13 TOIT; 0, [TMB], = 1, [O,] = 14 Torn; V, [TMB], = 0.5 [O,] = 14.5 Torr; x , [TMB], = 0.25, [O,] = 14.75 Torr; 0, [TMB], = 0.1, [O,] = 14.9 Torr.R. R. BALDWIN, M. w. M. HISHAM, A. KEEN AND R.w. WALKER 1173 Integration of eqn (xx) for small consumptions gives eqn (xxi): A[IBE]/t = 2k, [TMB] + 2k, k , [TMB] [MI t. (xxi) Defining kobs = A[IBE]/2[TMB] t kobs = k1+ k1 k,[M] t. The intercept of the plot kobs = A[IBE]/2[TMB] t against t should thus give kzbs = k,, the value when the OH chain is eliminated. Fig. 2 shows that the plots of kobs against z are linear. However, the intercept varies systematically with [TMB]. This arises because of the residual HO, chain, so that the rate equation is given, not by the simple expression (i), but by expression (iii). Thus, kzbs is given by eqn (xxii): kzbs = k, + k, (k,/k,)+ [TMBII. (xxii) A plot of the extrapolated values of kEbs from fig. 2 plotted against [TMBIf gives a reasonably linear relationship, from which k , = 2.7 x s-l, k,/k$ = 0.52 (dm3 mol-l s-l)+ at 440 OC (fig.3). However, the extrapolation of the kobs against t plot is sensitive to experimental error in measuring the small yields of IBE in the early stages of reaction, since at other temperatures the lines through the kobs against t points do not always show the expected increase in gradient with increasing [TMB]. Greater accuracy in the value of k , and particularly of k,/k$ can be obtained by using a computer treatment which, in effect, puts a mean line of the correct gradient for the particular mixture through each set of (kobs, t) points. Numerical integration of the differential eqn (iv)-(ix) is ([TMB 1 /Torr)j FIG. 3.-Plot of k& against [TMBf. 440 O C , data from intercepts of fig. 2.1174 DE C 0 M PO S I TI 0 N OF 2,2,3,3-TE T R A M E THY L B U T A NE carried out as for KCl-coated vessels, except that a differential equation is always written for H,O,.Six TMB + 0, mixtures were chosen, with concentrations ranging from 0.1 to 4.0 Torr at a total pressure of 15 Torr, except at the lowest temperature of 400 O C where 30 Torr of 0, was used with the total pressure adjusted to 60 Torr with N,. For each mixture, the yield of IBE was measured at four times where the consumption of TMB was c 5%; under these conditions secondary reactions of IBE are negligible. As with KCl-coated vessels, an optimisation procedure adjusted values of the parameters k , and k,/k$ so as to give minimum r.m.s. deviation between observed and calculated IBE yields. The only significant parameters determining the calculated [IBE] against t curve are k,, k,/k$ and the parameters for decomposition of H,O,.k, is accurately known, with M = H,, but the value of M,, given by eqn (xxiii), depends on the value taken for TMB relative to H, in reaction (8). Because of the low pressure of 0, and N,, TMB can make major contributions to M, [the value of M in reaction (S)]: M, = 0.35[0,] +O.43W2] + a(TMB) [TMB]. (xxiii) At each temperature, optimum values of k, and k,/k$ were obtained for values of a(TMB) = 1, 2.5, 5.0 and 10 relative to H, = 1. (The values of 0.35 and 0.43 for 0, and N, have been obtained in previous studies?) Unfortunately, the minimum r.m.s. 0.6 X x \ x “E I W Y P .E? i . 0 - n n n L 1.24 1.28 1.32 1.36 1.40 1.44 1.&8 103 KIT FIG.4.-Plot of log k,/k+ against 1/T. Aged B,O,coated vessels: x , u(TMB) = 1 ; 0, a(TMB) = 2.5; 0, u(TMB) = 5; A, u(TMB) = 10. KCl-coated vessels: 0, a(TMB) = 2.5. deviation for each value of a(TMB) did not vary sufficiently to eiable the optimum value of a(TMB) to be determined. However, increasing a(TMB) had relatively little effect on k,, whereas the value of k,/k$ decreased significantly, as shown in fig. 4. Fig. 4 also gives the values obtained for k,/k,f over the range 41 8-466 O C in KCl-coated vessels. In KCl-coated vessels, where a value of a(TMB) = 2.5 was used, the value for k,/k$ is fairly insensitive to a(TMB), partly because the contribution from H,O, dissociation is smaller, and partly because higher pressures of 0, and N, were used, so that the contribution of TMB to M, is less important; variation of a(TMB) from 1 to 5 had negligible effect on k,, and only altered k,/k,f by 1%.Fig. 4 shows that inconsistent variation of k,/kf with temperature is obtained when a(TMB) is givenR. R. BALDWIN, M. w. M. HISHAM, A. KEEN AND R. w. WALKER 1175 the unlikely high value of 10, and that a value of 2.5 gives the best consistency with the results in KCl-coated vessels. The best line through all the results, both in B203- and KC1-coated vessels, gives A,/A7i = 4.4 x lo5 (dm3 mol-l s-')!, E4 -4E7 = 81.7 f 8 kJ mol-l. If the present value15 of k, = 2 x lo9 dm3 mol-1 s-l, independent of temperature, is accepted, A , = 1.97 x 1Olo dm3 mol-1 s-l, E4 = 81.7f 8 kJ mol-l. The value obtained for k , is insensitive to a(TMB), the variation being an increase of 20% at 4OOOC and an increase of 4% at 52OOC as a(TMB) is increased from 1 to 10.The results, together with the r.m.s. deviation and the corresponding values of k , obtained in KCl-coated vessels, are summarised in table 4. TABLE 4.-vALUES OF k , IN B,03- AND KCI-COATED VESSELS _____~ ~ aged B,03coated vessel KC1-coated vessel r.m.s. dev. T/"C k,/s-l (%) k,/s-' - 1.45 x 15.8 400 420 6.17 x 5.7 440 2.77 x 10-5 3.6 2.56 x 10-5" 470 1.70 x 10-4 3.5 2.01 x 10-4" 500 1.37 x 10-3 3.7 1.33 x 10-3" 520 4.05 x 10-3 3.8 3.98 x 10-3a - 1.29 x - - 542 a These values differ slightly from those quoted in ref. (2) because of allowance for the interaction between homogeneous and heterogeneous termination. The high r.m.s. deviation at 400 OC reflects the very small amount of reaction at the lowest TMB concentrations even with reaction times of 2000 s.Measurements at 400 and 420 OC were included to examine the possible effect of an inefficient surface destruction of H,O,. However, both at 400 and 420 OC, introduction of values of k , in the range 0.001-0.008 s-l increased rather than reduced the values of the r.m.s. deviation. The values obtained for k , only changed by ca. 2% over the range k, = 0-0.008 s-l, whereas k4/k7i changed by ca. 15%. Little error is thus likely in the values of k , and k,/k7i due to surface destruction of H202. The best straight line through the values of k , in B203-coated vessels (omitting the value at 470 "C) gives A , = 1.26 x 1017 s-l, El = 295.8 kJ mol-l. Combination of the results for the B203- and KC1-coated vessels in table 4 gives A , = 1.04 x 1017 s-l, El = 294.7 kJ mol-l.These values may be compared with A , = 1.08 x 1017 s-l, El = 295.1 kJ mol-l obtained from the studies in a KC1-coated vessel. The differences are within experimental error. Examination of the scatter of the [IBE] against time data and the log k , against l/Tdata suggests that individual values of k , are accurate to within f. 4%. A 4% error in k, at the extremes of the temperature range would give an error of 1 % in El. The estimated accuracy of El is thus within + 3 kJ mol-l, consistent with the statistical accuracy of ca. 1%. The very close agreement with previously published values2 ( A , = 1.20 x 1017 s-l, El = 295.4 kJ mol-lj confirms the validity of the suggested thermodynamic data2q l6 for the t-butyl radical.It has been suggested', that the decomposition of H20, on a KCl-coated surface1176 DECOMPOSITION OF 2,2,3,3-TETRAMETHY LBUTANE may produce HO, radicals, although the fraction of H,O, giving HO, radicals is not stated. The agreement of the values of k, with the two vessel coatings indicates that the previously published value of El, from which theenthalpy and entropy of the t-C,H, radical have been calculated, has not been distorted by such production. Moreover, the agreement of the two sets of values of k,/k,g suggests that in the range 420-470 OC such production of HO, radicals does not significantly increase the concentration of HO, in KCl-coated vessels. This work was supported by the United States Office of Scientific Research through the European Office of Aerospace Research, United States Air Force. A grant from the S.R.C. for provision of gas-chromatographic facilities is gratefully acknowledged. G. A. Evans and R. W. Walker, J. Chem. SOC., Faraday Trans. 1, 1979, 75, 1458. G. M. Atri, R. R. Baldwin, G. A. Evans and R. W. Walker, J. Chem. SOC., Faraday Trans. 1, 1978, 74, 366. R. R. Baldwin, I. A. Pickering and R. W. Walker, J. Chem. SOC., Faraday Trans. I , 1980, 76, 2374. R. R. Baldwin and L. Mayor, 7th Int. Symp. Combustion (Butterworths, London, 1958), p. 8. R. R. Baldwin and D. Brattan, 8th Znt. Symp. Combustion (Williams and Wilkins, Baltimore, 1962), R. R. Baldwin, D. Jackson, R. W. Walker and S. J. Webster, Trans. Faraday SOC., 1967, 63, 1676. R. R. Baldwin and J. Howarth, J. Chem. SOC., Faraday Trans. I , 1982, 78, 451. R. R. Baldwin, A. R. Fuller, D. Longthorn and R. W. Walker, J. Chem. SOC., Faraday Trans. 1,1974, 70, 1257. R. R. Baldwin, M. J. Matchan and R. W. Walker, Trans. Faraday SOC., 1971, 67, 3521. York, 1967). 68, 1362. p. 110. lo J. 0. Hirschfelder, C. F. Curtiss and R. B. Bird, Molecular Theory of Gases and Liquids (Wiley, New l1 R. R. Baldwin, A. R. Fuller, D. Longthorn and R. W. Walker, J. Chem. SOC., Faraday Trans. 1, 1972, l2 D. R. Blackmore, J . Chem. SOC., Faraday Trans. 1, 1978, 74, 765. l 3 R. R. Baldwin, A. Keen and R. W. Walker, unpublished work. l4 A. B. Nalbandyan, 17th Int. Symp. Combustion (The Combustion Institute, Pittsburgh, 1978), p. 533. l5 R. W. Walker, in Reaction Kinetics (Spec. Period. Rep., The Chemical Society, London, 1975), l6 R. R. Baldwin, R. W. Walker and Robert W. Walker, J. Chem. SOC., Faraday Trans. 1, 1980,76,825. vol. 1, p. 161. (PAPER 1/799)

ISSN:0300-9599    

DOI:10.1039/F19827801165

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1982, 78, 1177-1187 Effect of Temperature on the Ionisation Constants of 2-, 3- and 4-Nitrobenzoic, Phthalic and Nicotinic Acids in Aqueous Solution BY LESLEY A. ASHTON AND JOSEPH I. BULLOCK* Cecil Davies Laboratory, Department of Chemistry, University of Surrey, Guildford GU2 5XH Received 18th May, 1981 The association constants (equivalent to pKd on the molar scale for the equilibria between hydrogen ion and the 2-, 3- and 4-nitrobenzoate, phthalate, hydrogen phthalate and nicotinate anions and the nicotinic acid zwitterion have been determined in aqueous solution at constant ionic strength from spectrophotometricdata at various temperatures between 288 and, in some cases, 473 K. 3- and 4-nitrobenzoic acids were stable to 473 K, 2-nitrobenzoic acid decomposed at 423 K and nicotinic acid at 408 K, and the absorption spectrum for the phthalic acid system was unsuitable for use above 448 K. Only one reaction, namely the protonation of the nicotinate anion [C,H,NO,]-, was isoelectric and the association constant describing the formation of the zwitterion decreased with increasing temperature.The plot of pK, against reciprocal temperature was linear within experimental error leading to temperature-invariant values of AH (exothermic) and A S with AC, = 0 for the association reaction. To 408 K, the association constant for the protonation of the zwitterion changed little in terms of the experimental errors. For 2-nitrobenzoic acid the plot of pK, against reciprocal temperature was also approximately linear. In this case the abnormally high acid strength at room temperature is caused largely by the highly endothermic nature of the association reaction and the acid becomes appreciably weaker at high temperature.For the rest the association constants (pK,) were greater at the highest temperatures employed than at 298 K which was largely caused by the high, positive entropy change for the association reaction resulting from a decrease in the dielectric constant of the bulk solvent as the temperature increased. Plots of pK, against reciprocal temperature were markedly non-linear with minima observed in some cases. These plots were analysed in terms of continually varying values of AH, A S and AC, over the temperature range although in some cases the changes in AC, were hardly significant in terms of the likely experimental errors.Previous from this laboratory has employed spectrophotometric data to study the protonation of heterocyclic nitrogen bases and, in some cases, the stability constants of iron@) complexes of the bases. The upper temperature limit was usually 473 K except where decomposition at a lower temperature occurred. An extension is to study the effect of temperature on the equilibria between the anions of organic acids and hydrogen ion. The computer program SQUAD^.^ allows the determination of association constants from spectrophotometric data alone and does not require predetermined extinction coefficients or the free hydrogen ion activity, and is thus particularly suitable for use above 373 K. Polyfunctional species can also be studied4 even if stages of the association reactions are overlapping across wide pH ranges.For ion-consuming reactions, e.g. the association of an acid anion or a zwitterion with a hydrogen ion, plots of pK, (log,, association constant) against 1/T are expected' to be non-linear and many show minima at relatively low temperatures. Continually varying values of A H and A S are expected with AC, interpreted as having a constant, non-zero value8 or varying over the temperature range9 or not reported at all1* depending on the precision of the experimental results. On the other hand, isoelectric reactions, e.g. the protonation of a neutral or the dissociation of 11771178 EFFECT OF TEMPERATURE ON IONISATION CONSTANTS a zwitterion, may give near-linear plots of ply, against 1/T with the extent of association decreasing with temperature in either case.EXPERIMENTAL The acids were used as supplied or recrystallised from water depending on the analytical purity of the materials supplied. The following analytical figures refer to the acids used in our experiments. Found: 2-nitrobenzoic acid, C 50.1, H 2.8, N 8.3%; 3-nitrobenzoic acid, C 50.4, H 2.9, N 8.4%; 4-nitrobenzoic acid, C 50.2, H 2.9, N 8.1 %; C,H,N04 requires: C 50.3, H 3.0, N 8.4%. Phthalic acid, found: C 57.4, H 3.4%; C,H,04 requires: C 57.8, H 3.6%. Nicotinic acid, found: C 58.5, H 4.0, N 11.5%; C,H,NO, requires: C 58.5, H 4.1, N 11.4%. Other materials were as described previo~sly.~ The spectrophotometers, high- and low-temperature cells, fumace4p l1 and the computer program SQUAD were used as b e f ~ r e .~ A summary of experimental conditions is given in table 1. The upper temperature limits were : 473 K, 3- and 4-nitrobenzoic acids, set4 by the experimental design; 423 K, 2-nitrobenzoic acid TABLE 1 .--SUMMARY OF EXPERIMENTAL CONDITIONS total total organic ionisable compounda no. of no. of range r acid hydrogen ion no. solutions wavelengths /nm /mol dm-3 mol dm-3 /mol dm-3 I 7 8 240-310 0.01 0.743 5.328 x 7.427 x I1 7 8 250-320 0.01 1.301 9.646 x I11 7 7 250-310 0.01 1.499 9.536 x 1.499 x IV 9 11 270-290 0.02 5.994 1.697 x 2.028 x V 10 7 244-268 0.05 2.046 1.553 x 1.301 x 10-4 1.044 x 10-4 a I, 4-nitrobenzoic acid; 11, 3-nitrobenzoic acid; 111, 2-nitrobenzoic acid; IV, phthalic acid; V, nicotinic acid.and 408 K, nicotinic acid, decomposition noted by observation of absorbance drift with time at higher temperatures; 448 K, phthalic acid, observed absorption envelope narrowed and flattened above this temperature making the spectrum unsuitable for use. The ionic strengths (table 1) were adjusted to constant values with AnalaR sodium chloride. With the nitrobenzoic acids series of solutions at fixed organic acid concentration but with varying hydrogen ion concentration (adjusted with standardised, AnalaR hydrochloric acid) were prepared. The absorbance of each solution was measured at various wavelengths at each temperature (table 1) and allowance made for the absorbance of the water-filled cell, which was large at high temperature because the cell windows are made from ~apphire.~ The procedures for phthalic and nicotinic acids were the same except that standardised, AnalaR sodium hydroxide solution was added to some solutions to generate high concentrations of the phthalate dianion or nicotinate anion.It was necessary to use more solutions and to measure them at a greater number of wavelengths for the dibasic than for the monobasic acids as the association reactions overlap over an appreciable part of the total acid concentration employed. The SQUAD input data consisted of the absorbance-wavelength matrix, the total organic acid concentration, the total conckntration of ionisable hydrogen ion, the expected species and an estimate of the equilibrium constant(s). No other information is required and a set of absorbance data can be readily obtained from a set of solutions prepared within the reagent conceqtrations outlined in table 1.L. A.ASHTON AND J. I. BULLOCK 1179 RESULTS AND DISCUSSION As b e f ~ r e l - ~ ~ l ~ the molar concentration scale was used throughout and it was assumed that the solutions expanded as would pure water.2*13 The density of water at each temperature was incorporated into the computer program. Constants were calculated for the association reactions : (i) for the nitrobenzoic acids, H+ + L- + HL loglOKF = pK, (ii) for phthalic acid, H+ + L2- HL- log,,K~ = pK,, and 2H+ + L2- e H2L log,,#F (iii) for nicotinic acid, H+ + L-" HL* log,,K~ = pK,, and 2H+ + L- e H2L+ logl,$? from which where HL* is the zwitterion. These values at the stated ionic strengths are listed in table 2.The uncertainties in parentheses were derived from the degree of fit between the experimental and computed spectra, and the true experimental errors must be greater than these. The fit was best for 4-nitrobenzoic acid because the observed absorption envelope (table 1) was well-separated from the tail of a much more intense short-wavelength absorption and for 3- and 2-nitrobenzoic acids the envelope shifted progressively to shorter wavelength so that for 2-nitrobenzoic acid there was appreciable mixing of the short- and longer-wavelength absorptions. TABLE ACIDITY CONSTANTS FOR THE ACIDS AT VARIOUS TEMPERATURES (a) 4-nitrobenzoic acid, (b) 3-nitrobenzoic acid (c) 2-nitrobenzoic acid, Z = 0.01 mol dm-3 Z = 0.01 mol dm-a Z = 0.01 mol dm-s 288 298 312 323 348 373 398 423 435 448 460 473 3.32 (1) 3.41 288 3.31 (1) 3.40 298 3.31 (1) 3.41 305 3.33 (0) 3.43 323 3.38 (1) 3.48 348 3.47(1) 3.58 373 3.55 (1) 3.67 398 3.68 (2) 3.82 423 3.75(2) 3.89 435 3.80 (7) 3.95 460 3.86 (3) 4.02 473 4.01 (1) 4.18 - 3.44(1) 3.53 288 3.40(2) 3.49 293 3.39(1) 3.48 298 3.42 (0) 3.52 305 3.43 (2) 3.53 323 3.48 (5) 3.59 335 3.50(5) 3.62 348 3.62 (4) 3.75 360 3.76 (1) 3.90 368 3.84(5) 4.00 373 3.99 ( 5 ) 4.16 398 - 423 - 1.93 (5) 2.02 2.05 (6) 2.14 2.16 (2) 2.25 2.28 (3) 2.38 2.39 (3) 2.49 2.55 (6) 2.65 2.61 (4) 2.71 2.67(1) 2.78 2.72 ( 5 ) 2.83 2.85 (6) 2.97 3.00(12) 3.13 2.11 (1) 2.201180 EFFECT OF TEMPERATURE O N IONISATION CONSTANTS TABLE 2.-continued (d) phthalic acid, I = 0.02 moldm-3 temp/K log,,KP PGl P G 288 298 305 323 348 373 398 423 435 448 460 4.65 (5) 4.69 (5) 4.65 (5) 4.76 (6) 4.91 (8) 5.15 (11) 5.25 (5) 5.41 (8) 5.50 (10) 6.56 (29) 5.55 (8) 7.48 (3) 7.50 (8) 7.40 (5) 7.64 (6) 7.83 (6) 8.10 (9) 8.36 (4) 8.59 (5) 8.70 (8) 8.77 (7) 9.85 (28) 5.13 5.17 5.14 5.28 5.46 5.75 5.90 6.13 6.25 6.35 2.95 2.93 2.87 3.00 3.05 3.09 3.27 3.36 3.38 3.42 (e) nicotinic acid, Z = 0.05 mol dm-3 288 298 303 318 333 348 363 378 393 408 4.78 (1) 4.72 (2) 4.68 (1) 4.63 (2) 4.57 (1) 4.47 (2) 4.39 (3) 4.24 (10) 4.14 (5) 4.10 (8) 6.90 (5) 6.71 (7) 6.67 (4) 6.60 (6) 6.58 (5) 6.45 (4) 6.35 (5) 6.24 (1 7) 6.17 (13) 6.22 (19) 4.96 4.88 4.85 4.80 4.75 4.66 4.58 4.45 4.36 4.33 2.12 1.99 1.99 1.97 2.01 1.98 1.96 2.00 2.03 2.12 For both phthalic and nicotinic acids it is possible to determine the equilibrium constants for the successive protonation reactions separately by working at very low or very high pH.However, as the association constants are temperature-dependent, at each temperature it would be necessary to check that the total acidity employed was sufficient to effectively remove the unwanted form. Using the program SQUAD this is unnecessary and the same total acidities and therefore the same set of solutions can be used for the entire temperature range. This simplifies the experimental procedures and is particularly useful above 373 K. Log KF and log /3F are correlated, in the computational method and for this reason overlapping pK, values for polybasic acids will always be less well-defined than those for monobasic acids assuming other factors to be comparable.In addition, the absorption envelope for phthalic acid was relatively narrow and small absorbance changes were observed. Above 448 K, the fit became very poor (table 2). For nicotinic acid, three forms are found across the acidity range, namely the zwitterion, [C,H,NO,] * , the protonated cation, [C,H,NO,]+, and the deprotonated anion, [C,H,NO,]-. At 298 K all three species coexisted to an appreciable extent in five of the ten experimental solutions. As with phthalic acid the observed absorption envelope was narrow and only small absorption changes were noted. Mean ionic activity coefficients calculated from the Davies equation1* (substituting the appropriate Debye-Huckel coefficient at temperatures other than 298 K) wereL. A.ASHTON A N D J. I. BULLOCK 1181 applied with the activity coefficients of the neutral or zwitterionic forms put equal to unity to give the so-called thermodynamic constants pKL shown in table 2. No correction was made to pKa2 for nicotinic acid on the assumptions that the activity coefficient of the zwitterion15 is equal to unity and that of the nicotinate cation is equal to that of the hydrogen ion at a given ionic strength. The contributions of the ions derived from the organic acids to the mean ionic strength were calculated using approximate pK, values at each temperature. The differences in ionic strength found when the refined values of pK, were used had a negligible effect. It was not possible to vary the ionic strength over a wide range by changing the concentration of the organic acids because the concentration required is largely determined by the extinction coefficients of the acid and its anion.However, by adding an inert electrolyte (in our case sodium chloride) it was possible to increase the ionic strength. This was done for 4-nitrobenzoic acid at 298 and 423 K (table 3). The values of pK, decreased with increasing ionic strength as expected, and on applying Davies activity coefficients14 there was a reasonable constancy when the possible sources of error were considered. TABLE EFFECT OF IONIC STRENGTH ON THE DISSOCIATION OF 4-NITROBENZOIC ACID 298 K 423 K I/mol dm-3 0.01 0.05 0.10 0.50 0.01 0.05 0.10 0.50 log,oKF 3.31 3.30 3.18 3.17 3.68 3.64 3.55 3.41 PK', 3.40 3.47 3.39 3.43 3.83 3.92 3.90 3.85 The values of pKi for all the acids compare reasonably well with the critical values of Martell and Smithls for 298 K and I = 0.These are: 4-nitrobenzoic acid 3.442 f 0.001 3-nitrobenzoic acid 3.449 0.001 2-nitrobenzoic acid 2.179 f 0.006 phthalic acid pK;, 5.408 pKL2 2.950 nicotinic acid pKL, 4.81 0.03 pKa2 2.05 _+ 0.03 The quoted deviations referle to agreement between different workers rather than to the significance of an individual measurement. Of the nitrobenzoic acids the 2-nitro-dekivative is much the strongest at 298 K and as the temperature is raised above ca. 320 K all of them become appreciably weaker. To 423 K, pK; for 2-nitrobenzoic acid decreases by an order of magnitude whereas for 3- and 4-nitrobenzoic acids the decreases to that temperature are much smaller.As the temperature is raised presumably the increasingly poor solvation of the proton more nearly offsets the steric hindrance in 2-nitrobenzoic acid. Phthalic acid (pKi,) is also comparatively strong. In the solid state" both carboxylic acid groups are inclined at 45.8O to the plane of the benzene ring, but in the acid anion1* (pH,]+ salt) the -COOH group is inclined at 26' and the -COO- group at 74O to the ring. As might be expected, pK;, for phthalic acid decreases with increasing temperature to a much greater extent than does pK,,. The second stage of ionisation of phthalic acid1182 EFFECT OF TEMPERATURE O N IONISATION CONSTANTS 4.2 4 .O lo3 KIT FIG. 1.-Variation of pG with 1/T. Curve A, 4-nitrobenzoic acid, left-hand scale; curve B, 3-nitrobenzoic acid, right-hand scale.(pK,,) is highly ionic-strength dependent.1° Lumme and KarilO found pK,, = 5.358 (298 K) from a least-squares treatment of the experimental pK, values and two expressions for activity coefficients with Z in the range 0.2668-2.0150 mol dm-3. At Z = 0.0570 mol dm-3 and 298 K they found pK,, = 4.963 which is higher than our value at Z = 0.02 mol dm-3. DERIVED THERMODYNAMIC QUANTITIES These were determined by the temperature variation method (fig. 1-3). As stated earlier, the protonation of the nicotinate anion is an isoelectric reaction on the assumption of the formation of the zwitterion. Earlier workla supports this and our plot of pK,, against 1 / T was linear. Temperature-invariant values of AH and A S were calculated (table 4) with the assumption that ACp = 0.The second stage! i.e. the protonation of the zwitterion, is non-isoelectric so that a non-linear plot of pKaz against 1/T might be expected. However, the equilibrium constant (table 2) changed little with temperature in terms of the degree of fit. This meant that AH, x 0 in terms of the experimental error. According to Christensen and co-workersZo AH, is ionic-strength independent in the range Z = 0.01 -0.09 mol dm-3 at 298 K which would seem to support our view that corrections to pKaz for activity coefficients need not be made. The critical valuesls for nicotinic acid are: 298 K, Z = 0, AH in kJ mol-l, A S in J K-l mol-l AH1=-11.7&1; AS1=5O AH, = - 3.3; There is good agreement with respect to AH, and AS, (table 3). ASz = 29.L.A. ASHTON AND J. I. BULLOCK 1183 5 . 0 1 4 I 1 I I I 1 I I 2.3 2.7 3 1 3.5 lo3 KIT FIG. 2. 3.0 2 . 8 - m % 2.6 2.4 2 . 2 2 .o I 1 I I 1 I I 2.4 2.8 3.2 3.6 103 KIT FIG. 3. FIG. 2.-Variation of pkf, with 1/T for phthalic acid. Curve (a) p&; curve (b) pkf,,. FIG. 3.-Variation of p G with 1/T for 2-nitrobenzoic acid: 0, present work; 0, Schaller, ref. (24). For 2-nitrobenzoic acid a plot of pKa against 1/T (fig. 3) was approximately linear so that a temperature-invariant value of AH was calculated from a least-squares treatment of this plot on the assumption that ACp = 0. AH and the calculated, constant value of AS are reported in table 4. Figures in' parentheses are derived from the standard deviation of fit for the pKa against 1/T plot. The enthalpy change is markedly positive for the association reaction and is the main cause of the abnormally high acid strength [see table 4, parts (u)-(c)].The solid-state structure of the acid reveals that the planes containing the carboxylic acid and nitro-groups are both inclined2' at high angles with respect to the plane of the benzene ring (24.1O and 54.3O, respectively) compared with the angles for the 3-nitro- (means 3O and 13S0, respectively22) and 4-nitro- (1.6' and 13.8O, re~pectively~~) benzoic acids. Presumably steric hindrance is reduced in the 2-nitrobenzoate anion although the structure is unknown. Whilst it is always possible to analyse data in terms of a constant, non-zero value of ACp [ref. (8) quotes 99.2 J K-l mol-'1 the possible sources of error hardly justify this in our case.The critical values16 for 2-nitrobenzoic acid are: 298 K, I = 0, AH = 14.1 kJ mol-l, AS = 89.1 J K-l mol-l, so that only fair agreement was found. These were determined by Everett and Wynne-Jonesa based on conductivity data of Schalleri4 reported in 1898. In order to determine AGO means of pKa at various dilutions were takenS8 However, the data24 show clear evidence of small but regular increases in pKa with dilution in the 2-nitrobenzoic acid range 7.81 x 10-3-9.77 x mol dm-3. For the most dilute solution pKa (298 K) = 2.19 and pKa (372 K) = 2.81; we found that the1184 EFFECT OF TEMPERATURE ON IONISATION CONSTANTS TABLE 4.-sELECTED THERMODYNAMIC QUANTITIES FOR THE ACIDSa* ' A H A S ACP temp/K /kJ mol-l /J K-l mol-l /J K-l mol-l 288 298b 312 323 348 373 398 423 435 448 460 473 288 298b 305 323 348 373 398 423 435 460 473 (a) 4-nitrobenzoic acid - 2.2 58 1.1 69 2.8 74 6.7 86 10.9 98 15.4 109 20.2 121 22.6 126 25.3 132 27.8 138 30.7 144 (b) 3-nitrobenzoic acid - 5.7 48 -3.1 57 -0.1 66 4.2 79 8.9 92 13.9 105 19.3 118 22.0 124 27.8 137 30.9 144 -0.8 (1 .O) 63 (3) -4.2 (1.1) 53 (9) 134 139 (1 1) 145 151 162 174 186 197 203 209 214 221 (27) 150 155 (43) 159 168 181 194 207 220 (61) 226 239 246 (c) 2-nitrobenzoic acid A H = 18.4 (0.5) kJ mol-l; A S = 103 (6) J K-l mol-l (d) phthalic acid 288 298b 305 323 348 373 398 423 435 448 8.3 9.3 (2.0) 10.0 11.9 14.7 17.7 20.9 24.3 26.0 27.9 ~ 119 123 (19) 125 131 139 148 156 164 168 173L. A.ASHTON AND J. I. BULLOCK 1185 TABLE 4.-continued (di) phthalic acid temp/K /kJ mol-l /J K-l mol-l AH2 AS2 288 29gb 305 323 348 373 398 423 435 448 2.6 3.3 (2.0) 3.8 5.2 7.2 9.4 11.8 14.3 15.6 17.0 65 67 (19) 69 73 79 85 91 98 100 104 (e) nicotinic acid (i) AH, = - 11.8 (0.7) kJ mol-l; ASl = 54 (3) J K-l mol-I (ii) 298 K: AH, x 0; AS, x 38 J K-l mol-l a Calculations based on pK', with AG = - 2.3026 RT pK', except AH, and AS, for nicotinic acid.Estimated errors in parentheses (see text). plot of PKa against reciprocal temperature was approximately linear. Our (pKa) results and those of S ~ h a l l e r ~ ~ are plotted together in fig. 3. Schaller did not employ solutions at constant ionic strength, but at the lowest acid concentLation the activity coefficients must be close to unity. For the rest, plots of pKa and (pK,) against reciprocal temperature (fig.1 and 2) were non-linear and were fitted to the polynomials shown below: 4-nitrobenzoic acid 2.3026 R ply: = 2.150 x lo4 T'-76.41+0.2331 T 3-nitrobenzoic acid phthalic acid 2.3026 R pKL = 2.725 x lo4 T'- 101.7+0.2600 T 2.3026 R pKal = 5.437 x lo3 T1+23.69+0.1662 T 2.3026 R pKa2 = 7.534 x lo3 T1-5.681 +0.1220 T from which values of AH, AS and ACp can be readily calculatedz5 at each temperature. The error in ACp is particularly large using this procedure25 with values of pKa derived from spectrophotometric data. In the present work, pKa values are best defined for 4- and 3-nitrobenzoic acids (see above) and ACp as well as AH and AS at each temperature are included in table 4. The standard deviations in these quantities may be estimated using a procedure outlined by King25 for equally spaced temperature intervals.Previous application^^^ used a mean temperature at or close to 298 K, where the deviations in A H and AS would be least, and therefore a very limited temperature range. In our experiments the mean temperature was much higher than 298 K so that at 298 K the calculated standard deviations in A H and AS are at their greatest; the deviation in ACp would increase with increasing temperature. Another disadvantage is that the method does not allow the low-temperature data, which were the most 39 FAR 11186 EFFECT OF TEMPERATURE O N IONISATION CONSTANTS precise, to be favourably weighted in the analysis. The method was applied to eight temperatures for 4-nitrobenzoic acid with pKa f 0.02 and six temperatures for 3-nitrobenzoic acid (pKa & 0.04) separated by 25 K intervals to give the results shown in parentheses in table 4.For phthalic acid (pKi&-O.10) seven temperatures were available but A H and A S (table 4) are less well defined because of the greater uncertainties introduced in pKa. Values of ACp are not included. The critical valuesl8 are : 298 K, Z = 0, A H in kJ mol-l, A S in J K-l mol-l; 4-nitrobenzoic acid A H = - 1.80f 1.3, A S = 59.8 3-nitrobenzoic acid A H = - 1.55 0.2, A S = 61.1 AH, = 2.68+0.04, AS, = 65.3. phthalic acid AHl = 2.09+0.17, AS, = 110 In general, the critical values were obtained from more precise measurements of PKa than ours but from very much smaller temperature ranges. For 3-nitrobenzoic acid A H was obtained directly26 from calorimetry but polynomial fits for pKL and T were also used2' as was the case28 for 4-nitrobenzoic acid.For phthalic acid it was foundlo that pKa = a+ bT. There is reasonably good agreement between our results and the critical values at 298 K except for A H for 3-nitrobenzoic acid and AH, for phthalic acid. The disagreement for phthalic acid arises from our apparently low value of pXal. A H is endothermic for 2-nitrobenzoic acid and for the rest becomes more endothermic for the association reaction as the temperature increases so that the decrease in acid strength is controlled by changes in A S which is always positive and becomes increasingly so as the temperature is raised (constant, positive A S for 2-nitrobenzoic acid). These findings are in keeping with the crude suggesting that AG is in a linear relationship with 1/e (where E is the bulk dielectric constant of AG = A+B/E water) leading26 to a proportionality between AG and A S and that the increase in A S is mainly a result of environmental effects largely concerned with the increasingly poor solvation of the proton as the temperature increases.At temperatures above 300 K plots of AG against l/e were approximately linear for 4- and 3-nitrobenzoic acids but curvature was observed above 400 K for both stages of ionisation of phthalic acid. In the range 298-473 K, E The uncertainty in the absorbance measurements is smaller below 373 K than above because of the reliabilities of the spectrophotometers used. For 4-nitrobenzoic acid, which was the best defined experimentally, the pKa values obtained from the SP 3000 used at low temperatures (288-348 K) were fitted to a polynomial of the type shown above and the equation used to predict pK, at 473 K. A value of 4.01 was obtained which was the same as the experimental quantity. Our results suggest that precise determinations of pKa over a limited, low-temperature range ( e g .278-333 K) may well be useful in predicting acidities to temperatures at least 100 K greater. For example, the best available data28 for 4-nitrobenzoic acid were fitted to give the results in table 5. At each temperature these results are indistinguishable in terms of the likely experimental errors. 4-Nitrobenzoic acid has28 a minimum value of pKa near 300 K and the data from ref. (28) consist of six determinations in the range 288-313 K separated by 5 K temperature intervals.For phthalic acid, the best literature values1* consist of four measurements in the range 288-308 K but in this case plots of pKa against T were linear for both stages 4 from 78.36 to 34.59.L. A. ASHTON A N D J. I. BULLOCK TABLE 5.-mDICTION OF VALUES FOR 4-NITROBENZOIC ACID 1187 AH AS T/K PK', /kJ mol -l /J K-l mol-l 473 4.14b 4.18' 28Sb 30.7' 139.7b 144' 298 3.441" 3.41' -0.51" -0.8' 64.2" 63' " Ref. (27); predicted from low-temperature data of ref. (27); ' this work. of ionisation. In the absence of a minimum value for pKi or marked curvature in the low-temperature range it is unlikely that the data will be of much value in pre- dicting high-temperature results.Financial support from the Central Electricity Generating Board and the S.R.C. is gratefully acknowledged. R. D. Alexander, D. H. Buisson, A. W. L. Dudeney and R. J. Irving, J. Chem. Soc., Faraday Trans. I , 1978, 74, 1081. D. H. Buisson and R. J. Irving, J. Chem. Soc., Faraday Trans. I , 1977, 73, 157. R. D. Alexander, A. W. L. Dudeney and R. J. Irving, J. Chem. SOC., Faraday Trans. 1,1978,74,1075. J. I. Bullock and P. W. G. Simpson, J. Chem. Soc., Faraday Trans. I , 1981, 77, 1993. D. J. Leggett and W. A. E. McBryde, Talanta, 1975,22, 781. Acetic acid, for example, H. S. Harned and R. W. Ehlers, J. Am. Chem. Soc., 1933, 55, 652. D. H. Everett and W. F. K. Wynne-Jones, Trans. Faraday Soc., 1939, 35, 1380. (I D. J. Leggett and W. A. E. McBryde, Anal. Chem., 1975, 47, 1065. @ F. S. Feates and D. J. G. Ives, J. Chem. Soc., 1956, 2798. lo P. Lumme and E. Kari, Acta Chem. Scand., Ser. A, 1975,29, 1 17. l1 R. D. Alexander, A. W. L. Dudeney and R. J. Irving, J. Phys. E, 1980, 13, 22. l2 A. S. Quist and W. L. Marshall, J. Phys. Chem., 1968, 72, 684. l3 G. C. Kennedy, W. L. Knight and W. T. Holser, Am. J. Sci., 1958, 256, 590. l4 C. W. Davies, Ion Association (Butterworths, London, 1962), p. 41. B. P. Kelley and T. H. Lilley, J. Chem. Soc., Faraday Trans. I , 1978, 74, 2779. l6 R. M. Smith and A. E. Martell, Critical Stability Constants (Plenum Press, London, 1976), vol. 1 and 3. l7 W. Nowacki and H. Jaggi, Z. Kristallogr., 1957, 109, 272. R. A. Smith, Acta Crystallogr., Sect. B, 1975, 31, 2508. R. W. Green and H. K. Tong, J. Am. Chem. Soc., 1956,78,4896. 2o J. J. Christensen, R. M. Izatt, D. P. Wrathall and L. D. Hansen, J. Chem. SOC. A, 1969, 1212. 21 S. S. Tavale and L. M. Pant, Acta Crystallogr., Sect. B, 1973, 29, 2979. 22 N. N. Dhaneshwar, S. S. Tavale and L. M . Pant, Acta Crystallogr., Sect. B, 1974, 30, 583. 23 S. S. Tavale and L. M. Pant, Acta Crystallogr., Sect. B, 1971, 27, 1479. 24 R. Schaller, Z. Phys. Chem., 1898, 25, 497. 2b E. J. King, Aci6Base Equilibria, International Encyclopedia of Physical Chemistry and Chemical Physics (Pergamon, Oxford, 1965), topic 15, vol. 4, chap. 8. T. Matsui, H. C. KO and L. G. Hepler, Can. J. Chem., 1974, 52, 2906. 27 P. D. Bolton, K. A. Fleming and F. M . Hall, J. Am. Chem. Soc., 1972,94, 1033. 28 J. M. Wilson, N. E. Gore, J. E. Sawbridge and F. Cardenas-Cruz, J. Chem. Soc. B, 1967. 852. 2@ (a) B. B. Owen, R. C. Miller, C. E. Milner and H. L. Cogan, J. Phys. Chem., 1961, 65, 2065; (b) G. C. Akerlof and H. I. Oshry, J. Am. Chem. Soc., 1950,72, 2844. (PAPER 1/800) 39-2

ISSN:0300-9599    

DOI:10.1039/F19827801177

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1982, 78, 1189-1 198 Threshold Energy and Excitation Function for the Reaction of Atomic Hydrogen with Cyclohexane BY DEREK GRIEF AND GEOFFREY A. OLDERSHAW* Department of Chemistry, University of Hull, Hull HU6 7RX Received 27th May, 1981 The reaction of photochemically generated hydrogen atoms with [aH,Jcyclohexane has been examined and the integral reaction probability of the abstraction reaction (1) determined at different translational (1) energies of H*: The phenomenological threshold energy of reaction (1) is 39 If: 4 kJ mol-l. Measurements of the moderating effect of xenon at different intial translational energies of H* have been combined with calculated collision densities to obtain the excitation function for reaction (1) in the energy range 30-200 kJ mol-l.The maximum reaction cross-section per D atom in this range is markedly less than that for the corresponding reaction of H* with the secondary D in n-C,D,,. H* + C6Dl, --* HD + C,D,,. Abstraction reactions of the hydrogen atom with alkanes have been examined by a variety of methods. In systems in which both reactants are essentially in thermal equilibrium with the surroundings, the Arrhenius parameters for several reactions have been determined from measurements of rate coefficients.l12 There is also literature concerning both the abstraction and substitution reactions of ‘hot’ tritium atoms, generated by nuclear recoil, with alkane^.^ Recoil tritium has initial energy far in excess of the maximum for chemical reaction and the reaction products result from the sampling of the excitation functions for different processes over the whole of their ranges.Only the most general information about reaction cross-sections can be obtained from such studies. Photodissociation provides an alternatiLe source of ‘hot’ H or D atoms which are reactive towards alkanes. The use of this method has the advantage that by varying the photolysis wavelength the initial translational energy of the atom can be selected in the range 0.3-4 eV. Following the original work on the reaction between D and H2,4 control of the initial energy has been exploited in several determinations of reaction Other work with alkanes has included the variation of reaction yields with initial energy of the atoms5* and examination of isotope effects.* However, the expectation that measurements of reaction yield as a function of initial energy would be used to obtain the energy-dependent form of the reaction cross-section (excitation function) has been fulfilled in only one instance, the reaction of H with butane.O The difficulty of calculating energy transfer in collisions between H and the polyatomic substrate was circumvented in that case by measuring the effect of initial energy on the moderating efficiency of xenon.The determination of the reaction cross-section then requires a knowledge of the collision density of H in xenon, which can be calculated from the H-Xe potential.1° We have now examined the reaction of photochemically generated H atoms with [2H,z]cyclohexane. Product yields for abstraction of D and the moderating efficiency of xenon have been measured for a series of initial atomic energies.Collision densities of H in Xe have been computed for each of the source energies and combined with 11891190 REACTION OF H* WITH [2H12]CYCLOHEXANE the experimental measurements to obtain the excitation function for the abstraction reaction. EXPERIMENTAL Gaseous mixtures of cyclo-C,D,, and HI or HBr were prepared in quartz vessels and irradiated with monochromatic light. In some experiments xenon was added to the reaction mixtures. Total pressures were in the range 20-500 Torr* and the temperature was ca. 293 K. H, and HD in the reaction products were determined mass-spectrometrically after condensing the reactants at 77 K. The mass spectrometer was calibrated with synthetic mixtures of H, and HD.For experiments with HBr the irradiation sources were a zinc resonance lamp, a cadmium resonance lamp and a super-pressure mercury lamp (Osram type HBO 200) with monochromator. The zinc lamp was used with a 100 mm filter of cis-but-2-ene at 100 Torr and the principal absorbed line was at 213.9 nm. The cadmium lamp was used with a 20 mm filter of toluene (1.0 x mol dm-3) in hexane; in this case the principal absorbed lines were at 228.8 and 226.5 nm with relative intensities of cu. 17: 1, giving an effective mean absorbed output wavelength of 228.7 nm. The mercury lamp was used with a monochromator setting to give a mean output wavelength of 248.4nm. For experiments with HI, the mercury lamp and monochromator were used to obtain mean output wavelengths of 312.4, 333.4 and 352.0 nm.In order to eliminate, at the higher wavelengths used, small quantities of low-wavelength radiation which would be preferentially absorbed, plate-glass filters of thickness 4 and 12 mm were used at the wavelengths 333.4 and 352.0nm, respectively. The half-width of the monochromator output was between 3 and 6 nm for the various wavelength settings. For irradiation at the highest wavelength, a combination of the super-pressure lamp, an interference filter and a 25 mm plate-glass filter was used to obtain an output wavelength of 365+3 nm. [ZH,Jcyclohexane (99%) was from Fluorochem, xenon from B.O.C. and Air Products, and HD (98%) used for calibration from Merck, Sharpe and Dohme. RESULTS AND DISCUSSION Photolysis of mixtures of HI and cyclo-C,D,,, or HBr and cyclo-C,D12, yielded products which included H, and HD. In the reaction system, H atoms with high translational energy are generated by photolysis of HX and may abstract D from cyclo-C,D,, to generate HD.Energy loss in non-reactive collisions also occurs and those atoms with energy below the threshold for reaction with cyclo-C,D,, are scavenged by HX to yield H,. Product ratios for various photolysis wavelengths are shown in fig. 1-3. For runs in which no xenon was present in the reaction mixture the ratio [H,]/[HD] was found to be a linear function of the reactant ratio [HX]/[C,D,,], with a positive intercept. As in other systems of this type5* the results are interpreted using the following reaction mechanism : H*+HX+H+HX (4) H*+Xe+H+Xe ( 5 ) H+HX+H2+X X+X+M +X2+M.* 1 Tom = (101 325/760)Pa.D. GRIEF AND G. A. OLDERSHAW 1191 I 1 I I I I 1 0 0.2 0.4 0.6 0.8 1.0 1.2 FIG. 1.-Product ratios in the photolysis of mixtures of HBr and cyclo-CeDl, at 229 nm. tHBr1 /tC,Di,] 0 0.2 0.4 0.6 0.8 1.0 1.2 FIG. 2.-Product ratios in the photolysis of mixtures of HBr and cyclo-CeDl, at 248 nm (0) and of mixtures of HI and Cyclo-CeD,, at 3 12 nm (a). [HXl/[CciDnI 80 - 60 - 5 - 40 z - n N Y 20 - I 1 1 1 I I 1 0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 FIG. 3.-Product ratios in the photolysis of mixtures of HI and Cyclo-CeD,, at 333 (0) and 352 (0) nm. [HII/[C,D,,l1192 REACTION OF H* WITH [2Hl,]CYCLOHEXANE H* respresents an atom with energy above the threshold for reaction (1) and H represents one with less than this energy.Steps (2), (4) and (5) are moderation processes and (4) includes both non-reactive collisions and moderation through the exchange reaction.ll9 l2 Abstraction by X of D from C,D12 or X from HX is excluded on energetic grounds. The hot substitution reaction (8) H* + C,D1, 3 C,DllH + D (8) has a threshold of ca. 1.5 eV139 l4 and is expected to be much less important than the abstraction (1) even at the highest energy (2.0 V) used in the present work. Thus for small extents of reaction the ratio of products is and in the absence of xenon [H,]/[HD] is a linear function of [HX]/[C,D,,]. The coefficients k, to k, depend on the initial energy of H* as reflected in the slopes and intercepts of fig. 1-3. Reaction of thermalised H atoms with accumulated X, from reaction (7) competes with reaction (6) as the photolysis proceeds.For runs with HI the consequent reduction in [H,]/[HD] was in the region of 2% and was ignored, but for runs with HBr the observed product ratios were increased by between 3 and 10% to correct for scavenging by Br,.15 INTEGRAL REACTION PROBABILITIES AND THRESHOLD FOR D ABSTRACTION The integral reaction probability (net probability of reaction) for a hydrogen atom of specified initial energy is P = kl/(kl + k2). Values of P obtained from the intercepts k 2 / k , of plots of [H2]/[HD] against [HX]/[C,D,,] are given in table 1. Experiments at 365 nm enabled an upper limit of 0.003 to be set on the fraction of HD in the reaction products for an initial energy of 35 kJ mol-l.The initial laboratory energy of the H atom EL was ca1culatedls*l7 using values of the dissociation energy Do of TABLE INTEGRAL REACTION PROBABILITY AND kl/k5 FOR DIFFERENT INITIAL ENERGIES OF THE HYDROGEN ATOM EL (initial) /kJ mol-l P k l / k 5 197 0.205 k 0.006 4.93 & 0.07 161 0.200 k 0.008 4.15 & 0.07 120 0.193 k 0.005 2.83 k 0.07 90 0.155 & 0.008 2.00 & 0.09 66 0.086 & 0.009 1.22 0.08 47 0.024k 0.003 - 35 - - HBr and HI of 362.5 and 294.7 kJ mol-l, respectively. Fig. 4 shows the variation of the integral probability of reaction (1) with the initial laboratory energy of the hydrogen atom. The phenomenological threshold obtained by linear extrapolation of the lowest points is EL = 39 f 4 kJ mol-l. As pointed out else~here,~~ l5 this does not represent the true threshold for reaction.Conversion to relative energy requires aD. GRIEF AND G. A. OLDERSHAW 0.05 1193 - / I & I I 1 0 50 100 150 200 EL/kJ mol-’ FIG. 4.-Integral reaction probability for reaction (1) at various initial laboratory energies of H. reduction of only 1 %, but model calculations using an assumed ‘line-of-centres’ excitation function shows that linear extrapolation of yields gives an apparent threshold substantially above the true value. Few other thresholds for abstraction of D by H are available for comparison. Gann et al.9 have determined the true threshold for abstraction of the secondary D in [2Hl,]butane as 34 & 2 kJ mol-1 and the phenomenological threshold determined by direct extrapolation of their data is ca.37 kJ mol-l. Since the secondary CH bonds in butane and cyclohexane have approximately the same strength1* the similarity of the thresholds for secondary D abstraction is not unexpected; for the stronger primary bond in CD,CH,CH,CD, Gann et aL9 give the apparent threshold for D abstraction as 48 f 5 kJ mol-l. Fink and coworkers have determined thresholds for abstraction by D of H from a number of alkanes.s* 19-21 In comparing their results with the present work it should be noted that the translational energies of D appear to be based on a value for the dissociation energy for HI slightly higher than that adopted here. For the purposes of comparison with the present work the thresholds given by Fink and coworkers should be raised by ca. 3 kJ mol-1 to bring the results onto a common basis.Allowing for this adjustment, the threshold reported21 for the abstraction by D of H from cyclohexane is equal to that found here for the isotopically reversed process. This is surprising in view of the difference in the CH and CD bond strengths, but the expected difference in thresholds is within the combined experimental error. MODERATING EFFECT OF XENON Addition of inert gases to the reaction mixtures reduced the yield of hot product HD by removing energy from the hot atoms in non-reactive collisions. A detailed study of the moderating effect of xenon was carried out principally to provide information for use in the evaluation of the excitation function, as discussed more fully below. If a comparison of the ratio [H,]/[HD] is made for two reaction mixtures, each having the same ratio [HX]/[C,D12], one with and the other without xenon moderator, then eqn (9) shows that the increase in [H,]/[HD] caused by the presence of xenon is1194 REACTION OF H* WITH [2H12]CYCLOHEXANE 0 10 20 30 40 50 60 FIG.6.-Moderating effect of xenon at 229 nm. [XeI/[C,D,21 0 10 20 30 40 FIG. 7.-Moderating effect of xenon at 248 nm. W e 1 /[C6D121 Experiments on the moderating effect of Xe were carried out using a number of different initial energies of the hydrogen atom. In each case measurements of ([H2]/[HD])mod were combined with values of ([H,]/[HD]),,,,, for the appropriate ratios [HX]/[C,D,,] obtained from fig. 1-3. The resulting values of A([H,J/[HD]) for the photolysis wavelengths used are shown in fig.5-8 and are proportional toD. GRIEF AND G. A. OLDERSHAW 1195 0 10 20 30 LO [Xel/[C$,,l FIG. 8.-Moderating effect of xenon at 312 (0) and 333 (0) nm. ~e]/[C,D,,] as required by eqn (10). Values of kl/k5 derived from the plots are listed in table 1. EVALUATION OF REACTION CROSS-SECTION The procedure is that suggested by Gann et aL99 l6 The reactivity of H* with some particular initial (source) distribution of kinetic energies (designated by subscript a) expressed relative to moderation by xenon is given by where Erepresents the relative translational energy of H and C8DI2, E, is the threshold (relative energy) for reaction (l), EL is the laboratory translational energy of H and SR ( E ) is the cross-section for reaction (1). S(E,),, is the average total scattering cross-section of xenon for a hydrogen atom of energy EL and n, (EL)Xe is the collision density in pure xenon of H atoms with initial E,-distribution a.G(EL,E) is the normalised distribution function of laboratory translational energies corresponding to a particular value of the relative translational energy E between H and C&2, the latter having a Maxwellian velocity di~tributi0n.l~ Values of kl/k5 are determined experimentally for different initial energies of the hydrogen atom by measuring the moderating effect of xenon, as described in the last section. By comparing the kl/k5 values for two different initial H atom energy distributions, a and b, the average reaction cross-section over the intermediate energy range can be obtained from where In addition to the experimental values of kl/k6, the collision densities of H in xenon are required and are computed from the H-Xe interaction potential.1196 REACTION OF H* WITH [2H12]CYCLOHEXANE COLLISION DENSITY OF H IN Xe Collision densities of the hydrogen atom in a thermal xenon medium were calculated using random selection of the dynamical variables in a manner similar to that outlined by Rebick and Dubrin.lo Hydrogen atoms of initial energy corresponding to that generated in the photolysis are imagined to make successive collisions with xenon atoms in thermal equilibrium at the experimental temperature.The collision density is obtained by computing the average number of collisions occurring in each energy interval. The computations are conveniently carried out using a H-Xe interaction potential of the form V(r) = drd and the Born-Mayer potential given by Bickes et aZ.22 was fitted to this form at 0.35 and 2.0eV to yield S = 9 and d = 0.329 J mol-1 nm9.In each collision between the hydrogen atom and the xenon atom the energy of the hydrogen atom after collision, EL, was calculated from that before collision, E L usingl5? 23 ,u EL = EL - [2EL - m,, vXe2 +(mXe-mH) (2EL/mH)avXecosyl -cOsx) +,u ( 2 ~ ~ / m ~ ) ~ v , , s i n y s i n ~ c o s q (14) where M = mH +mxe, ,u is the reduced mass of H and Xe, vXe is the xenon speed, y is the angle between the H and Xe velocity vectors and x (deflection angle) and q (azimuthal angle) are the scattering angles [as shown in ref. (lo)]. Taking the vXe distribution to be Maxwellian at the experimental temperature, vXe, y and q were chosen randomly from suitably weighted distributions.l0? l5 For the inverse power repulsive potential x is a function of yo = b (E/8d)1/6,24 where b is the impact parameter, and the relationship between x and yo was computed for S = 9.15 x was determined by a random choice of yo, i.e.yo = yomaX&, where R is a random number between 0 and 1. To avoid counting collisions with very large impact parameters which involve small deflections and energy losses, an arbitrary constant value of yomax of 1.1 was employed over the whole energy range. This excluded collisions with energy losses less than ca. 1 % of the maximum. As a consequence of this procedure, the total scattering cross-section varied as E-&; In selecting the initial energy of the hot atoms a distribution appropriate to each photofysis wavelength was used.This took into account the monochromator band pass, the translational motion of HX and the distribution of HX among rotational states. For energies below ca. 93% of the source energy an asymptotic formulalo* 24 appropriate to the functional form of S(E),, adopted here was used to calculate the collision density : 4,uA(1)(9) (EL - 3kT/2)-' where A 9 9 ) = 0.327.24 DERIVED EXCITATION FUNCTION Int (a), Int (b), . . .were computed for each of the source energies from the collision densities na(EL)xe, nb(EL)Xe,. . .and used in eqn (12) with (kl/k5)a, (kl/k5)b,. . . to obtain the average reaction cross-section forD. GRIEF A N D G. A. OLDERSHAW 1197 141 12 / I I I I I 0 50 100 150 200 ElkJ mol-I FIG.9.-Excitation function for reaction (1). over the interval between each adjacent pair of source energies. The derived excitation function is shown in fig. 9. The lowest point shown was calculated from the experimental kl/k5 at EL = 66 kJ mol-1 and a value of EL = 47 kJ mol-1 obtained by extrapolation of a plot of k1/k5 against EL. Extrapolation to a lower energy was avoided because of the unknown shape of the kl/k5 against EL curve near the threshold. Owing to the fact that the lowest initial energy for which kl/k5 was determined (66 kJ mol-l) is much higher than the threshold, the value of the true threshold energy and the shape of the excitation function below 60 kJ mol-1 are difficult to establish. The function drawn in fig.9 in this region is of the ‘line-of-centres’ form, &(E) cc (1 - Eo/E),withEo = 30 kJ mol-’,butasteeperfunctionwithhigherthreshold is also compatible with the results. Taking into account both the values of kJk5 and the values of P listed in table 1, the true threshold lies in the range 28-38 kJ mol-l. The observations are of course insufficient to reveal details of the shape of the excitation function in the immediate region of the threshold. In the case of abstraction by H of secondary D from n-C4Dlo (16) Gann et al. found9 that the best fit to their results was obtained with an excitation function rising from the threshold significantly steeper than the line-of-centres function, but pointed out the uncertainties surrounding the shape of the function in the threshold region.The other principal feature of the excitation function determined for reaction (1 6)9 was the presence of a maximum in the cross-section at ca. 116 kJ mol-l. In the present case, the results obtained for abstraction by H of D from cyclo-C,D,,, reaction (l), allow the possibility of a maximum around 150 kJ mol-1 but do not establish one conclusively. In any event, the sharp decline in the excitation function over the range 100-180 kJ mol-l found for reaction (16) is not observed for reaction (1). The curve drawn in fig. 9 is of the line-of-centres form up to 140 kJ mol-1 and falls slightly below this function at higher energies. On the assumption that a maximum in the cross-section for reaction (1) exists below H* + n-C4Dlo + s-C,D, + HD1198 REACTION OF H* WITH [2H12]CYCLOHEXANE 200 kJ mol-l, it is substantially smaller, when allowance is made for the number of D atoms, than the maximum cross-section found for reaction (16).In the case of reaction (1) the maximum cross-section is (1 1.8 _+ 2.0) x 1 0-3 nm2, or (0.98k0.17) x nm2 per D. By using a number of different H-Xe potential functions Gann et aL9 obtained a maximum cross-section for reaction (16) of (3.5 f 1.5) x nm2. Treatment of their experimental data using the 9th power potential employed in this work yields a maximum cross-section of (6.3 f 1.2) x nm2, or (1.58 k0.30) x nm2 per D. Thus the strengths of the bonds broken, and the threshold energies, for reactions (1) and (16) are similar, but the excitation functions in the region below 200 kJ mol-l differ significantly.The rate coefficient of the related reaction D +CyC10-C6H12 + HD +~yclo-C~H,, has been measured as a function of temperatu~e.~~ In general, the activation energy is expected to be higher than the true threshold energy, which is itself lower than the phenomenological threshold. The activation energy reported for reaction (1 7), 16.7 & 1.3 kJ mol-l, is substantially lower than that expected from the observed threshold of reaction (17)21 or that of reaction (1). D.G. thanks the S.R.C. for a research studentship for the period when this work was carried out. W. E. Jones, S. D. MacKnight and L. Teng, Chem. Rev., 1973,73,407. R. R. Baldwin and R. W. Walker, J. Chem. SOC., Faraday Trans. 1, 1979, 75, 140. F. S . Rowland, in Chemical Kinetics, ed.J. C. Polanyi (Butterworths, London, 1972), p. 109. A. Kuppermann and J. M. White, J. Chem. Phys., 1966,44,4352. G. A. Oldershaw, in Gas Kinetics and Energy Transfer, ed. P. G. Ashmore and R. J. Donovan (Specialist Periodical Report, The Chemical Society, London, 1977), vol. 2, p. 96. P. Vidaud, R. D. Fink and J. E. Nicholas, J. Chem. SOC., Faraday Trans. I , 1979, 75, 1619. ' G. A. Oldershaw and E. A. Robinson, Chem. Phys. Lett., 1978,54, 527. ti G. D. Beverly and R. M. Martin, J. Phys. Chem., 1976,80, 2063. lo C. Rebick and J. Dubrin, J. Chem. Phys., 1970, 53, 2079. l1 D. J. Malcolme-Lawes, J. Chem. SOC., Faraday Trans. 2, 1978, 74, 182. l2 H. Y. Su, J. M. White, L. M. Raff and D. L. Thompson, J. Chem. Phys., 1975, 62, 1435. l3 C. C. Chou and F. S. Rowland, J. Chem. Phys., 1969,50, 2763. I4 M. Menzinger and R. Wolfgang, J. Chem. Phys., 1969, 50, 2991. l6 R. G. Gann, W. M. Ollison and J. Dubrin, J. Am. Chem. SOC., 1970,92,450. R. G. Gann, W. M. Ollison and J. Dubrin, J. Chem. Phys., 1971,54, 2304. D. Grief, Ph.D. Thesis (University of Hull, 1979). G. A. Oldershaw and D. A. Porter, J. Chem. SOC., Faraday Trans. I , 1974, 70, 1240. R. W. Walker, in Reaction Kinetics, ed. P. G. Ashmore (Specialist Periodical Report, The Chemical Society, London, 1975), vol. 1, p. 161. l8 J. E. Nicholas, F. Bayrakceken and R. D. Fink, J. Chem. Phys., 1972, 56, 1008. 2o F. Bayrakceken, P. Vidaud, R. D. Fink and J. E. Nicholas, J. Chem. SOC., Faraday Trans. I , 1976, 21 R. D. Fink and J. E. Nicholas, J. Chem. SOC., Faraday Trans. 1, 1972,68, 1706. 22 R. W. Bickes, B. Lantzsch, J. P. Toennies and K. Walaschewski, Faraday Discuss. Chem. Soc., 1973, 2s R. N. Porter, J. Chem. Phys., 1966,45, 2284. 24 J. 0. Hirschfelder, C. F. Curtiss and R. B. Bird, Molecular Theory of Gases and Liquids (Wiley, New 25 P. Kim and R. B. Timmons, Znt. J. Chem. Kinet., 1975, 7, 143. 72, 1058. 55, 167. York, 1964). (PAPER 1/853)

ISSN:0300-9599    

DOI:10.1039/F19827801189

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. Sac., Faraday Trans, I , 1982, 78, 1199-1207 Infrared Spectroscopic Study of the Effects of Different Cations on NN-Dimethylacetamide and Fully Deuterated NN-Dime t hylformamide BY W. EARLE WAGHORNE* AND HECTOR RUBALCAVA Department of Chemistry, University College Dublin, Dublin 4, Ireland Received 27th May, 1981 The infrared spectra of NN-dimethylacetamide and fully deuterated NN-dimethylformamide complexed to a variety of metal cations in propan-1-01 have been measured. The OC-N stretching and 0-C-N bending frequencies were found to vary systematically with the electrostatic potential at the surface of the cation. The 0-CN stretching frequency was primarily determined by the electronic structure of the complexed cation. The data indicate that complexation takes place via the carbonyl oxygen for all of the complexes studied, and support the idea of electrostatic stabilization of negative charge at the carbonyl oxygen by cations.The data also indicate that back bonding may take place between transition metal cations and complexed amides. In a recent p.m.r. study1 it was found that the barrier to rotation about the OC-N bond of NN-dimethylacetamide (DMA) complexed to different cations increased systematically with the charge and radius of the complexed cation. The silver ion, which lowered this rotational barrier, was exceptional among the cations studied. It has also been found that the lH and 13C chemical shifts of complexed amides2v vary systematically with the electrostatic potential arising from the complexed cations.Surprisingly, while amide complexes have been the subject of several infrared spectroscopic studie~,~-lO no such correlation with the nature of the complexed cation has been reported for the observed changes in any of the amide bands. Thus it was decided to carry out an infrared spectroscopic study of simple amides complexed to a wide range of metal cations. To avoid complications arising from the participation of water in the cation-amide interactions1l7l2 it was decided to carry out the work in non-aqueous media, and propan-1-01 (PrOH) was chosen as solvent. DMA and fully deuterated NN- dimethylformamide ([2H,]DMF) were chosen as amides for the study. EXPERIMENTAL Propan- 1-01 was dried over anhydrous CaSO, and fractionally distilled. DMA was purified by fractional distillation under reduced pressure as described previously;13 CCl, (Merck Spectroscopic grade) and [,H,]DMF (Fluorochem 99%) were used without further purification. The spectra of CCl, and propan-1-01 showed no significant impurity peaks. The perchlorates of barium, calcium, cobalt, nickel and strontium were prepared by reacting the carbonates with perchloric acid.The other electrolytes were analytical grade or better, except Zn(C10,), (Alpha, 98.9%) and Pb(ClO,), (Alpha, 95 %) and were not further purified. All of the electrolytes were dried at 60 OC under reduced pressure for 24 h and stored over P,O,. Subsequent analysis' indicated that AgClO,, LiClO, and MgCl, were anhydrous, Mg(C10,), existed as the monohydrate and Cd(C10,),, Cu(ClO,),, Pb(C10,), existed as the 11991200 INFRARED S T U D Y OF AMIDE COMPLEXES hexahydrates.In making up solutions it was assumed that the other electrolytes existed as the stable hydrates :I4 Al(ClO,), 6H,O, Ba(C10,), 3H,O, Ca(ClO,), 2H,O, Co(ClO,), 6H,O, Ni(C10,), 6H,O and that Sr(ClO,), was anhydrous. The molecular sieves (B.D.H., 4A) used to dry the electrolyte solutions were exposed to 1 mol d n r 3 solutions of the appropriate electrolyte until no further cationic exchange was detected, washed with distilled water and activated at 250 "C under reduced pressure for 12 h. Infrared spectra were measured using a Perkin-Elmer 125 spectrometer and matched liquid sample cells having a nominal path length of 0.1 mm and AgCl windows. The spectra were recorded from 4000 to 400 cm-', and the range from 2000 to 800 cm-l was recorded in an expanded mode.The experimental precision, for a sharp absorption band, was 2 cm-l for the former and 1 cm-l in the expanded range. Peak positions were measured relative to those of a sample of polystyrene. Solution spectra were measured using either the solvent or the appropriate M(C10,), solution as reference. The latter was preferred for the electrolyte solutions as it reduced the interference caused by the ClO; absorption bands, particularly in the region 1000 to 1150 ern-'. RESULTS The possibility of effects arising from the formation of amide dimers in solution7* was investigated by recording the spectra of DMA solutions in PrOH for the concentration range 0.1-1 .O mol dm-3. No change was observed in the frequencies of any of the amide absorbance bands, indicating that dimerization of the amides should not affect the results of the present study.The spectra of DMA solutions (0.5 mol dm-3) containing varying amounts of water (to 2 mol dm-3) were measured. The only significant change in the DMA spectrum was a broadening of the absorption band at 1635 cm-l. This could be removed by using PrOH containing the same concentration of water in the reference cell, indicating that the observed broadening resulted from a superposition of the water and amide bands and that the presence of water did not affect the amide bands in PrOH. The spectra of DMA and [,H,]DMF solutions containing AgClO,, LiClO,, Cd(ClO,),, Cu(ClO,),, Mg(ClO,),, Pb(ClO,), and Zn(ClO,), were recorded after they had been dried with the appropriately treated molecular sieves and after the addition of known amounts of water (to 1 mol dm-3).The frequencies of all the amide bands remained unchanged. Therefore it was concluded that the presence of small amounts of water in the solutions did not affect the measurements. The only differences between the spectra of DMA and [,H,]DMF solutions exposed to the various molecular sieves and those of the undried solutions was the presence of the water bands in the latter. Thus it is clear that, over the required drying time, neither the amides nor the PrOH were significantly decomposed by the molecular sieves. The spectra of DMA solutions (0.5 mol dm-3 DMA) containing varying concentra- tions of AgClO, and LiClO, (to 2 mol dm-3), Mg(ClO,), (to 1 mol dm-3), Pb(ClO,), and Zn(C10,), (to 0.5 mol dm-3) were recorded.The bands of the uncomplexed amide became progressively weaker and those of the complexed amide more intense with increasing electrolyte concentration. The positions of the bands associated with the complexed amide, where they were sufficiently removed from those of the uncomplexed amide to be clearly resolved, were independent of the electrolyte concentration (see fig. 1). In all cases the positions of the bands of the complexed amides were determinable for solutions having a cation to amide ratio of 1 to 1 [i.e. 0.5 mol dm-3 M(ClO,),, see fig. 11. The spectra of 0.5 mol dm-3 [2H7]DMF solutions containing 0.5 mol dm-3 Mg(C10,), and MgCl, differed only by the presence of the Cloy bands in the former.This indicated that the anion was not influencing the positions of the absorption bands of the complexed amide. Since the frequencies of the amide bands were not affected by changes in theW. E. WAGHORNE AND H. RUBALCAVA 1201 1700 1600 1700 1600 v/cm-' FIG. 1.-Variation of vco of DMA (0.5 mol dm-3) in the presence of varying LiC10, (I) and Pb(ClO,), (11): DMA ratios. I, [LiClO,]:[DMA] = (a) 0; (b) 0.5; (c) 1.0; ( d ) 2.0: 11, [Pb(ClO,),]:[DMA] = (a) 0; (b) 0.125; (c) 0.25; ( d ) 0.50; (e) 1.00. TABLE IN INFRARED FREQUENCIES OF DMA AND [2H,]DMF DMA [2H,]DMF neatb in CCldC in PrOHC neatb in CClaC in PrOHC 30 1 5dm 2925ds 1746m 1650%~ 1545w 1 500dm 1410s 1390s 1350 1260s 1 180s 1055s 1030w 1008w 953vw 730dw 585ds 469dm - 2930d 1742 1 660d 1546 1 494d 1410 1392 1352 1264 1184 1055 1030 1008 - 584d 463d - 1635d 1555 1 500d 1414 1400 1260 1185 - - 735d 590d 470d 22 1 Odm 2 1 30dm 2 1 05dm 2065dm 1 700ds 1 650evs 1390vs 1260s 1123s 1070m 1050m 1035m 918m 890s 835w 765dw 615ds 22 1 Od 2130d 2 1 OSd 2065d 1 69lId 1665,d 1644sh 1380, 1388sh 1264 1123 1070d 1050d 1032d 916 890 835 612d - 22 1 Od 2150d 2 1 05d 2070d 1 700d 1644d 1400 1260 1123 - 892 - 620d a Units are cm-'; uncertainties f 1 cm-l unless otherwise stated; w, m, s, vs, sh indicate weak, medium, strong, very strong and shoulder, respectively.* As their films between KBr discs. As 0.5 mol dm-3 solutions, 0.1 mm path length. f 2 cm-l. 5 cm-l.1202 INFRARED STUDY OF AMIDE COMPLEXES concentrations of the amide or electrolyte, by the presence of water nor by the nature of the anion present, it was concluded that the observed changes resulted from direct cation to amide interactions. Table 1 lists spectral data for DMA and [2H,]DMF as thin films formed between KBr discs and as 0.5 mol dm-3 solutions in CCl, and PrOH.Several of the amide bands were not observable in the solution spectra because of dilution or interference by the solvent bands. Tables 2 and 3 list the frequencies of the observable bands, below 2000cm-1, in the spectra of DMA and C2H,]DMF (0.5 moldm-3) in solutions containing 0.5 mol dm-3 of the different electrolytes. Several of the amide bands were obscured by those of CIO,. TABLE 2.-INFRARED FREQUENCIES OF DMA IN Mn+-DMA COMPLEXES IN PrOHa complexed cation ~~ Li+ 1 640b 1 50Sb 1418 1400 1264 596b Ag+ Mg2+ 1630b 1520 1418 1404 1260 cu2+ 1 604b 1511 1420 1402 1259 592b Zn2+ 1610b 1514b 1420 1402 1260 595b Cd2+ 1610b 1512b 1419 1401 1260 595b Pb2+ 1 596b 1508 1418 1401 1259 590b 1 608b 1 50gb 1414 1498 1260 - - a Measured for solutions containing 0.5 mol dm-3 DMA and 0.5 mol dm-3 M(ClO,),; units are cm-l; uncertainties f 1 cm-l unless otherwise noted.Values are f 2 cm-l. TABLE ~.-INFRAFED FREQUENCIES OF [2H,]DMF IN Mn+-[2H,]DMF COMPLEXES IN PrOHa complexed cation Li+ Ag+ Mg2+ Ca2+ Sr2+ Ba2+ co2+ Ni2+ CU2+ Zn2+ Cd2+ Pb2+ ~ 1 3 + 1642,b 1648sh 1 632b 1 6Ub 1 640b 1 640b 1 639,b 1630sh 1633b 1 632b 1 630b 1 632b 1 627b 1618,b 1612sh 1644,b 1650sh 1412 1404 1421 1418 1412 1408 1414 1415 1412 1414 1412 1409 1420 1260 1260 1251 1249 1250 1260 1250 1250 1245 1250 1248 1243 1240 893 894 900 898 897 896 900 900 902 900 900 900 904 630b 622b 64Sb 635b 628b 620b 645b 650b 666b 645b 640b 632b 685b a Measured for solutions containing 0.5 mol dm-3 [2H,]DMF and 0.5 mol dm-3 M(ClO,),; Values are units are cm-l; uncertainties f 2 cm-l.1 cm-l unless otherwise noted; sh, shoulder. DISCUSSION The simplest 'zero-order' description of the vibrations of an amide molecule, using non-interacting valence force coordinates, predicts several localized vibrations, including a carbonyl stretching, an OC-N stretching and an 0-C-N bending among others. Clearly a more accurate description would involve mixing of theseW. E. WAGHORNE A N D H. RUBALCAVA 1203 deformations to give the more realistic normal vibrational modes.Nevertheless, the simple zero-order description is sufficient for our present purposes, since only the major contributions to three of the amide absorption bands are required. Thus a complete vibrational assignment of the spectra of DMA and [2H7]DMF was not required nor was one attempted. However, it is necessary that the major contributions to the bands of interest are known before the effects of complexation can be understood. The strong bands at 1654 and 1635 cm-l in the spectra of [2H7]DMF and DMA, respectively, have been assigned to the amide carbonyl stretching vibration ( ~ ~ ~ ) , ~ - ~ 9 15* l6 although some participation by other deformations, most importantly the OC-N stretching vibration (vCN), has been suggested in the case of [2H7]DMF.15 The strong band at 1400 cm-l in the [2H,]DMF spectrum and the weaker one at 1500 cm-l in the DMA spectrum have been assigned to the OC-N stretching vibration (vCN).,-'9 1 5 9 l6 The moderately strong band at 620cm-l in the [2H7]DMF spectrum has been assigned to a vibration which is predominantly the 0-C-N bending vibration (docN) with some contribution from the rocking and stretching vibrations of the -N(C2HJ2 group.No corresponding band in the spectrum of DMA was observed. It is clear from the data in tables 2 and 3 that complexation affects the positions of the 1645 cm-l [2H7]DMF and 1635 cm-l DMA bands similarly, as it does the 1400 cm-l [2H7]DMF and 1500 cm-l DMA bands. Thus the assignment of these pairs of bands to similar vibrational modes is reasonable. None of the observed DMA bands showed variations similar to that of the 620 band of [2H,]DMF.It is necessary to establish, as far as possible, the predominant site of the cation to amide interactions before the effects of this interaction are discussed. The oxygen and nitrogen atoms of the amide molecule are the most likely sites because of their relatively high electronegativities. However, considerations of the amide structure in terms of simple resonance theory indicate that the oxygen should be the preferred site. Thus the bonding in an amide molecule can be described as a hybrid of structures I and II,17a \ / R R C=N+ R C-N R I \ R -0 I I1 each making a significant contribution. Clearly the resulting negative charge at the oxygen and positive charge at the nitrogen will strongly favour interactions at the oxygen atom.The coulombic effect of complexation at the oxygen should be to increase the relative importance of structure I1 to the description of the amide bonding. That is, the 0-CN bond order should decrease and that of the OC-N bond increase. Conversely, complexation at nitrogen should greatly decrease the relative importance of 11, thus reducing the OC-N bond order and increasing that of the 0-CN bond, relative to the unperturbed molecule. It is clear that, in the absence of complexation, the bonding in the amide molecule will be influenced by the amide-solvent interactions. For example, interactions with the hydroxyl proton of PrOH should have effects similar to those of complexation to a cation. However, in all cases, the vc0 frequencies of the complexed amides are lower and the vCN frequencies higher than those of the neat amides or of the amides dissolved in CCl,.This allows one to tentatively conclude that complexation occurs via the amide oxygen. / \ 01204 INFRARED STUDY OF AMIDE COMPLEXES The question of the site of complexation can be approached in a second independent way. It has been shown' that complexation of DMA to Na+, Li+, Mg2+, Cd2+, Pb2+ and Zn2+ raises the activation energy for rotation around the OC-N amide bond, indicating an increase in the OC-N bond order. This unambiguously places the interaction at the carbonyl oxygen for these complexes. It is clear from the preceding discussion that the infrared spectra of oxygen-complexed amides should differ markedly from those of nitrogen-complexed amides.However, there is a general consistency among the spectra of the amides complexed to the closed-shell cations, including Li+ and Mg2+, and among those of amides complexed to the transition-metal cations, including Cd2+ and Zn2+ (cf. below). Thus we can conclude that each of the cations is complexed via the carbonyl oxygen of the amides. This agrees with the conclusion of other w o r k e r ~ ~ ' ~ ~ that the amide oxygen is the predominant site of complexation to cations. 0 5.0 10.0 9/10 J C-' FIG. 2.-Variation in the frequency vCN for complexed [2H,]DMF with the electrostatic potential y at the surface of the complexed cation ; solid line, transition-metal cations ; broken line, closed-shell cations. Fig. 2 shows the variation in the frequency of the vCN band of [2H7]DMF with changing electrostatic potential, ry, at the surface of the complexed cation.The electrostatic potential was calculated via 2 1 ry=-X-- rc ~xE,, where 2 is the ionic charge, rc is the ionic radius 17b and E, is the permittivity of free space (E, = 8.85 x 10-l2 F m-l). As noted above the variation of vCN for DMA is similar to that for [2H,]DMF (cf. tables 2 and 3). Since ry is a purely electrostatic parameter of the complexed cation the increase in vCN with ry is in agreement with the simple electrostatic argument. However, the data in fig. 2 suggest that the nature of the complexed cation is also important. There is a linear variation (solid line) of vcN for [2H7]DMF complexed to those cations which have d electrons in their outer shell and a separate variation (broken line) for those which do not.At low values of ry this latter variation is much more pronounced than the former. It is interesting to compare the variations in vCN of the complexed amides with thoseW. E. WAGHORNE AND H. RUBALCAVA 1205 in the activation energy, AE,, for rotation around the OC-N bond of DMA complexed to different cati0ns.l The AE, values for DMA complexed to Li+, Cd2+, Mg2+, Pb2+ and Zn2+ were also found to increase with increasing ly,l and there is a good linear correlation between the vCN and AE, values for these DMA complexes. Since the AE, values refer unambiguously to the OC-N amide bond this correlation strongly supports the assignment of the vCN band. 680 660 I 5 51 --- z b o 640 r - - - 1 ($g%:; , , ., , , , 62 0 P B O 0 5.0 10.0 $/lo J C-' FIG. 3.-Variation in the frequency of 6,,, for complexed [2H,]DMF with the electrostatic potential y at the surface of the complexed cation. The AE, value of DMA complexed to Ag+ is unusual, in that it is less than that of the uncomplexed amide.lV 1 9 9 2o However, the vCN values of the amides complexed to Ag+ are not anomalous, which indicates that the predominant Ag+ complex involves the expected increase in the OC-N bond order. Thus the lower AE, value observed for DMA complexed to Ag+ does not result from an anomaly in the bonding of the predominant species in solution. This strongly supports an earlier suggestion19* 2o that the observed lowering of AE, by Ag+ results from the presence of a small but kinetically significant proportion of the amide molecules being complexed to Ag+ via the nitrogen atom.This is similar to the accepted mechanism for the lowering of AE, by the protonation of amides21 Fig. 3 shows the variation in the frequency of the dOCN band of [2H,]DMF with ly. There is a relatively good linear correlation for these data. However, the values for the Ba2+, Li+ and Sr2+ complexes lie below the line, which suggests that there could be separate correlations for complexes of the cations having available d electrons and those without, as in the case of the vCN data. The value of 666 cm-l for the Cu2+ complex is a striking anomaly which we cannot presently explain. Fig. 4 contains the vco data for complexed [2H,]DMF as a function of ly.There is no apparent correlation between these spectral data and ly. Rather, the frequency of the vco band appears to be related to the electronic structure of the complexed cation. Thus vco is 1642 _+ 4 cm-l for complexes with closed-shell cations, 1630 _+ 4 cm for complexes with the transition-metal cations and 1618 cm-l for the Pb2+ complex. The fact that the vco frequencies of the amides complexed to the transition-metal cations and Pb2+ are lower than those of the amides complexed to the closed-shell1206 INFRARED STUDY OF AMIDE COMPLEXES cations is particularly interesting. In each case the vco frequency is relatively independent of ly, indicating that interactions other than simple coulombic ones are involved. The lowering of vco by Pb2+ and the transition-metal cations is consistent with a degree of covalent bonding involving the metal d electrons and the amide carbonyl 1640 - I 1630- .ou - T T I 0 Cd 1621 OPb I T 1 OAl 1 1 0 5.0 10.0 $/lo J C-' FIG. 4.-Vanation in the frequency of vco for complexed [2H,]DMF with the electrostatic potential y at the surface of the complexed cation. n antibonding orbitals (back bonding).22 Such an interaction would weaken the amide 0-CN bond and result in the observed lowering of the vco frequency. This interaction could also give rise to the observed difference between the dependences of vCN on ly for the transition-metal and the closed-shell cations. We thank Mr G. Flynn for making the infrared measurements. W. E. Waghorne, A. J. I. Ward, T. G. Clune and B. G.Cox, J. Chem. SOC., Faraday Trans. I , 1980, 76, 1131. A. Fratiello, D. P. Miller and R. Schuster, Mol. Phys., 1967, 12, 111. Ch. P. Rao, P. Balaram and C. N. R. Rao, J. Chem. SOC., Furaday Trans. I , 1980, 76, 1008. W. E. Bull, S. K. Madan and J. E. Willis, Znorg. Chem., 1963, 2, 303. A. J. Carty, Can. J. Chem., 1966, 44, 1881. E. W. Randall, C. M. Silcox Yoder and J. J. Zuckermann, Inorg. Chem., 1966, 5, 2240. M-H. Baron, C. de Loze and G. Sagon, J. Chim. Phys., 1973,70, 1509. P. Combelas, M. Costes and C. Garrigou-Lagrange, Can. J. Chem., 1975, 53, 442. lo C. N. R. Rao, H. S. Randhawa, N. V. R. Reddy and D. Chakravorty, Spectrochim. Acta, Part A , 1975,31, 1283. l1 M. J. Adams, C. B. Baddiel, R. G. Jones and A. J. Matheson, J. Chem. SOC., Furaday Trans. 2, 1974, 70, 1114. M. J. Adams, C. B. Baddiel, G. E. Ellis, R. G. Jones and A. J. Matheson, J. Chem. SOC., Furaday Trans. 2, 1975, 71, 1823. ' M-H. Baron, J. Corset, C. de Loze and M. L. Josien, C.R. Acad. Sci., Ser. C, 1972, 274, 1321. l3 T. G. Clune, W. E. Waghorne and B. G. Cox, J. Chem. SOC., Faraday Trans. I , 1976, 72, 1294. l4 Handbook ofchemistry and Physics, ed. C. R. Weast (Chemical Rubber Co., Cleveland, Ohio, 1968).W. E. WAGHORNE A N D H. RUBALCAVA 1207 l5 G. Durgaprasad, D. N. Sathyanarayana and C. C. Patel, Bull. Chem. Soc. Jpn, 1971, 44, 316. l6 K. L. Dorris, T. H. Siddal, W. E. Stewart and M. L. Good, Spectrochim. Acta, Part A, 1967,23,1657. L. Pauling, The Nature ofthe Chemical Bond (Oxford University Press, London, 2nd edn, 1952), (a) p. 207, (b) p. 343. D. N. Waters, Z. Kantaret and N. N. Rhamna, J. Raman Spectrosc., 1978, 7 , 288. l9 P. A. Temussi and F. Quadrifoglio, Chem. Commun., 1968, 844. 2o P. A. Temussi, T. Tancredi and F. Quadrifoglio, J. Phys. Chem., 1969, 73, 4227. 21 B. G. Cox, J. Chem. Soc. B, 1970, 1780. 22 F. A. Cotton and G. Wiikinson, Advanced Inorganic Chemistry (Interscience, London, 1962). (PAPER 1/855)

ISSN:0300-9599    

DOI:10.1039/F19827801199

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1982, 78, 1209-1221 Diffusion of Dextran at Intermediate Concentrations BY BARRY N. PRESTON* AND WAYNE D. COMPER Department of Biochemistry, Monash University, Clayton, Victoria, Australia AND ANTHONY E. HUGHES, IAN SNOOK AND WILLIAM VAN MEGEN Department of Applied Physics, Royal Melbourne Institute of Technology, Melbourne, Australia Received 27th May, 1981 The diffusion properties of dextran molecules in water have been followed by use of boundary relaxation techniques (refractive index and tracer measurements) and by photon correlation spectroscopy. Measurements have been carried out up to concentrations of ca. 250 kg m-3. The diffusion coefficients and their inter-relationships have been interpreted in terms of the non-ideal behaviour of flexible polymers.The kinetic and equilibrium behaviour of macromolecules in a concentrated phase is of interest for many diverse aspects in biology and chemistry. Our studies1v2 in particular have been aimed at obtaining a better understanding of the dynamic behaviour of the extracellular matrix of connective tissues, which is known to contain high concentrations of polysaccharide chains. In this paper we present an analysis of the diffusional properties of dextran. We consider that the study of dextran polysaccharides serves as a useful, simplistic model by which the fundamental dynamic properties of concentrated polysaccharide regions in tissues may be understood. We have determined both the mutual diffusion coefficient and the intradiffusion coefficient of several dextran preparations of molecular weight varying from lo4 to 1.5 x lo5 by boundary relaxation techniques.These studies have been carried out over an extended concentration range (up to 250 kg m-3). For comparison, we have used photon correlation spectroscopy (P.c.s.), a non-perturbing technique, to investigate the motions of the dextran chains in similar solutions. Our experimental studies have been considered in relation to various theoretical treatments of diffusion and the semi-dilute properties of flexible p01ymers.~-~~ This work continues and expands earlier studies of the diffusion of macromolecules in concentrated aqueous solution^.^^ lo THEORY The transport of material by diffusion in a binary non-electrolyte system is usually described in terms of three coefficients which can be derived directly from experimental measurements.They are the differential mutual diffusion coefficient D, which is a measure of the rate of inter-diffusion of components 1 and 2, and the self-diffusion coefficients 0; and D,*. (We define component 1 as solute and component 2 as solvent.) It has been shown that the intradiffusion coefficient, D:, of labelled molecule 1* is at least approximately equal to the self-diffusion coefficient 0; of component 1,3 i.e. D l = 0;. (1) 12091210 DIFFUSION OF DEXTRAN Furthermore, the different types of diffusion are related by4 where c, is the molar concentration and p1 the chemical potential of component 1. This expression for D has been used by various investigators in polymer diffusion studies.6-10 Note that eqn (2) is used here with reservation, as it is based on the assumption of a regular s~lution,~ which is probably not the case for dextran solutions. An expression of the term (apl/ac,),, in terms of thermodynamic non-ideality coefficients can be derived from algebraic expressions for the chemical potential of component 1 given by Ogston5 as (3) where m, is the molality of component 1 [mol (g solvent)-l] and a,, a3... are the coefficients expressing thermodynamic non-ideality. Differentiation of eqn (3) with respect to m, and conversion of units of concentration into mass/volume units (C,) with eqn (2) gives p,-& = RT(lnm,+a,m,+a,m~+. . .) (1 +2A,M,C1+3A3M,q+. . .) (4) 0;' (1 - c, K) D = where and where is the partial specific volume of 1.This expression for D is different from that described previously by Yamakawa6 and used by various inve~tigators~-~~ in polymer diffusion studies. Yamakawa6 found that the (1 - C, F) term occupied the numerator instead of the denominator as in eqn (4). The difference is derived from the fact that Yamakawas assumed C2 = 1, as is approximated in dilute solutions, for his expression of osmotic pressure [his eqn (30.43)]. Furthermore, Yamakawa6 and others12 have evaluated fluxes relative to a cell-fixed frame of reference with the assumption that the volume change on mixing is negligible. In using the consistent expressions of chemical potential in terms of molal quantities as described by Ogston5 no assumptions are further required in the description of eqn (4) except for 5 being independent of C,.An alternative view of polymer dynamics in concentrated polymer solutions lies with the concept of an entangled statistical network mesh. This approach has been the subject of extensive theoretical work.l1* l3 At concentrations equal to or greater than the overlap concentration C* (noting that C = C* corresponds to a close-packed system of non-overlapping coils) divergence from dilute solution behaviour is anticipated. To calculate C* the polymer molecule is arbitrarily viewed as occupying either a sphere of radius R, or a cube of side R,, which gives rise to upper and lower M < C * < L . 3M1 4nNR; NR& bounds : ( 5 ) Further development by de Gennes13 has shown that the mutual diffusion coefficient (also termed a cooperative diffusion coefficient) in the semi-dilute region will vary with concentration as C0.75.The special case of flexible polymer movement within the statistical network mesh has been described by de Gennes as reptation and gives the intradiffusion coefficient 0: cc C-1.75. It is of interest to investigate to what degree the diffusional behaviour of dextran in water (a good solvent) is in agreement with these predictions.PRESTON, COMPER, HUGHES, SNOOK AND VAN MEGEN 121 1 EXPERIMENTAL MATERIALS The polymer dextran samples were either supplied or kindly donated by AB Pharmacia (Uppsala, Sweden). The physicochemical properties of the fractions used in this study are described in table 1. Note that the dextran fraction FDR7782 (with an HW/mn ratio of 1.32) was a purified subfraction obtained by gel chromatography of dextran T150.TABLE 1 .-PROPERTIES OF DEXTRAN SAMPLES virial coefficientsb moisture dex tran H" M W Mza content second (A,) third (A,) classification / 10" g mol-' / l(r g mol-I Hw/Hn / 10' g mol-' / 10-0 g(H,O) g-I /lo-!!? g-I / 10- mol ma g-* / 10-8 mol ma g-4 T10 0.6 1.04 1.68 1.08 9.89 8.7 0.7 2.4 T20 1.65 2.04 1.24 2.28 4.62 12.0 T70 3.95 6.95 1.76 7.36 7.22 23.0 0.47 2.0 T150 8.60 15.40 1.80 16.20 7.90 35.5 FDR7783 12.03 15.82 1.32 17.86 7.46 37.9 0.33 3.1 Calculated assuming symmetrical distribution Hz = Hn[3 -(2Mn/aw)]. From references in ref. (5) and (14) and unpublished results. PREPARATION OF LABELLED DEXTRANS A sample of dextran (5-50 mg) was dissolved in sodium hydroxide (0.01 mol dm-3, 0.5 cm3), [3H]KBH, solution (10 mCi ~ m - ~ , 10 mm3, Radiochemical Centre, Amersham, Bucks) was added and the mixture was stored at 50 OC overnight.The solution was acidified with glacial acetic acid to destroy untreated borohydride and neutralised with sodium hydroxide (1 mol dm-3). The polymer was separated from tritiated water and salts on a Sephadex G-25 column. All preparations were then dialysed extensively against water except for the dextran T10 preparation, which was purified by repeated ultrafiltration with water as solvent using an Amicon UM0.5 filter (Amicon Corp., Mass.) at 2.8 kg m-2. A comparative analysis of the 7 m 5 5 r5 E 1 K,V FIG. 1.4hromatographic profile of [3H]FDR7783 (0) and unlabelled FDR7783 (0) on Sephadex 6BC1 (column dimensions 90 cm x 1.5 cm) in H,O.The [SH]FDR7783 fractions were monitored for radioactivity; 97.5 % of the initial counts placed on column were recovered. The unlabelled FDR fractions were monitored for hexose;lS 100% recovery was obtained for this sample.1212 DIFFUSION OF DEXTRAN distribution of labelled and unlabelled dextrans by gel chromatography revealed no significant difference (fig. l), suggesting that the labelling procedure does not result in any significant change in the molecular size distribution of the dextran. PREPARATION OF DEXTRAN SOLUTIONS Stock solutions of dextrans of known moisture content were made up by weight in either doubly distilled water or in 0.15 mol dm-3 NaCl. Samples for photon correlation spectroscopy were then centrifuged for 1.5 h at CQ.100000 g before being transferred with sterile syringes into the light-scattering cells. All transfers were performed in a dust-free flow cabinet. DETERMINATION OF DIFFUSION COEFFICIENTS MUTUAL DIFFUSION COEFFICIENTS Mutual diffusion coefficients were measured by free diffusion in a Beckman model E analytical ultracentrifuge using the Schlieren optical system (this technique will be referred to as the refractive index method, abbreviated as rim.). A solution of concentration C, was layered over a solution of concentration C, (with C, > C,) in a synthetic-boundary cell and the speed was reduced to ca. 5000 r.p.m. The concentration step, AC = C,-C,, was close to 5 kg m-3 for routine measurements. Diffusion measurements for solutions of concentration 2 kg m-3 were made with C, = 0.Diffusion was observed at 20 O C in cells with 0.5, 1.2 or 3.0 cm synthetic-boundary centrepieces. The apparent diffusion coefficients, D,,?, were calculated from the broadening of the Schlieren peak by following the change in width at the half-height. If the diffusion is followed by measurement of the broadening of the peak at (dn/t)x),,,,, denoted by T, then 2 Dapp = ( ?Q2/ 16t In 2 where t is the time. Values of Dapp were calculated from the slope of the plot of ( W)z against t. The assignment of Dapp as a parameter describing some average diffusion coekcient for a FIG. 2.-The mutual diffusion coefficient of FDR7783 obtained by r i m . as a function of the difference in concentration across the boundary AC for the mean dextran Concentration maintained at a constant value of 35.5 kg m-3 (V) and 10.1 kg m-3 (A).The values represented by the curve (0) have been obtained by keeping the concentration of dextran C, constant at 35.5 kg m-3.PRESTON, COMPER, HUGHES, SNOOK AND VAN MEGEN 1213 polydisperse system (such as the dextrans used in this study) is regarded as an approximation only. Yet experimental evidence cited below demonstrates that the effect of molecular-weight polydispersity of the sample on its diffusion coefficient suggests that this factor is of minor significance. We shall therefore make the assumption that Dapp is identical to the mutual diffusion coefficient, D, at the mean concentration16 C[ = (C, + cb)/2]. We have performed a limited number of experiments to estimate the error associated with this assumption for a system exhibiting concentration-dependent diffusion coefficients.Fig. 2 shows the variation of DaPp with AC studied at two different constant values of C, namely 10.1 and 35.5 kg m-3 for dextran FDR7783. (Note that this dextran has low-order polydispersity with &fw/Hn = 1.32.) At both mean concentration values, Dapp values decrease linearly with increasing AC. This is in contrast with the predicted variation of Dapp with (AC)2 for monodisperse material, as given by Gosting and Fujita." While the variation of D,,, with AC is found, the value of Dapp at C = 5 kg m-3 is only ca. 3% less than the value obtained by linear extrapolation to AC = 0. In view of this relatively small error, Dapp values measurcd at AC = 5 kg mP3 will be regarded as the mutual diffusion coefficient at the corresponding C value.It is evident from the variation in Dapp with AC (fig. 2) that there is no identity between Dapp and D, the mutual coefficient, for measurements performed at non-zero AC for concentration-dependent systems. In a further attempt to explore an identity of Dapp with the initial concentration on either side of the boundary, experiments were performed with cb constant at a value of 35.5 kg m-3 while AC was varied. In this case, an even more marked, non-linear dependence of Dapp on AC was found, although the extrapolated value of Dapp at AC = 0 for this series gave a similar value of Dapp to those obtained from experiments performed with C constant at 35.5 kg m-3. INTRADIFFUSION COEFFICIENTS These were measured following Park's modification18 of the open-ended capillary method.D* lD All runs were conducted in a water bath at 20f0.01 O C .At the end of the run, the capillary was removed, rinsed in water and the contents centrifuged directly into a vial for radioactive counting. The radioactivity was measured by standard counting techniques in a Packard Tri-Carb model 3003 liquid-scintillation spectrometer. No attempt was made to correct these diffusion coefficients for boundary phenomena, as suggested by Nanis et as the diffusion coefficients obtained by this technique were identical, within experimental error, to those obtained by other techniques.12 DIFFUSION COEFFICIENTS FROM PHOTON CORRELATION SPECTROSCOPY A Spectra Physics argon-ion laser (type 2600, A,, = 488.0 nm) was focused onto a scattering cell in a bath at ambient temperature (20 f 0.50 "C).Light-scattering fluctuations were detected by a Malvern precision device photomultiplier assembly type S/N 101 112 (with a D2608 E.M.I. photomultiplier tube). The output signal was correlated by a Malvern 72 channel digital correlator type K7023 operated in the scaling mode. The field autocorrelation function g(l) (K, z) of the Kth spatial fourier component is related to the diffusion coefficient D by the equation21 g(1) (K, z) = C-flDt. This treatment assumes that the solute is monodisperse, in whch case a plot of In [Ig(l)(K, z)l] against Kzz yields a straight line of slope -D. If the system is polydisperse there is no general theory relating molecular parameters to g(l) (K, t) for concentrated solutions.In this paper the correlation function has mainly been analysed by the method of cumulants. We have found that for the dextran concentrations studied the quadratic cumulant used suitably describes the experimental 23 (fig. 3). No depolarised scattering was observed for the samples studied, indicating that multiple scattering was insignificant. RESULTS A N D DISCUSSION DIFFUSION COEFFICIENTS AT INFINITE DILUTION (Do) The values of diffusion coefficients at infinite dilution were obtained by manual extrapolation of diffusion coefficients obtained at non-zero concentrations (outlined1214 DIFFUSION OF DEXTRAN 0 10 20 30 40 50 channel number FIG. 3.-(a) Variation of the field, autocorrelation function with channel number (sample time 1.85 ps per channel) for dextran FDR7783 at a concentration of 28.0 kg m-3.Experimental data (a), quadratic cumulant fit (-). (b) As (a) with dextran FDR7783 at a concentration of 132 kg m-3. below), as we have no prior knowledge of the manner of their concentration dependence. The results obtained from the three techniques used, i.e. r i m . , P.C.S. and intradiffusion, are presented in table 2. The error associated with measurements of r.i.m. and P.C.S. was < 5%. However, owing to the relatively strong concentration dependence of 0: values at low concentrations (outlined below), the error in the (D +)o values was near 10%. The Do values obtained by r i m . and P.C.S. (table 2) were found to be in excellent agreement. For the range of polydispersity encountered with these samples, the diffusion coefficients are relatively insensitive to the different weighted averages of Do that arise from r.i.m.and P.C.S. Since gel-chromatographic analysis did not reveal any clear indication of degradationPRESTON, COMPER, HUGHES, SNOOK AND VAN MEGEN 1215 TABLE 2.-DIFFUSION COEFFICIENTS AT INFINITE DILUTION diffusion coefficient (Do)/ 10-l' m2 s-l dex tran from classification from r.i.m. from P.C.S. intr diffusion T10 9.9 8.9 10.9 T20 6.7 6.7 T70 3.5 3.3 4.5 Tl5O 2 . 2 2.2 3.7 FDR7783 2.19 2.2 3.7 of the dextran following the labelling procedure, it is suggested that the difference between ( D + ) O and Do may be due to a non-uniform distribution of label in the preparation. Scaling the values of Do obtained from r i m .gives the relations DO = (1 -62 & 0.07) x 10-4j@;0-552( k0.005) DO = (1.15+0.50) x 10--4M-0.543(+0.045) and Do values from P.C.S. gives the relation D o = (1.17+0.3) x 10--4&f--0.521(k0.021). In the latter case, values of the exponent are found to be higher than that obtained from the P.C.S. data of Sellen24 on the diffusion coefficients of dextran in dilute solution, where he obtained an exponent value of 0.45. These differences are probably within experimental error, as his Do values are similar to those obtained in his study. The most significant difference between our data and Sellen's corresponds to measurements on the T10 sample. As pointed out by Sellen,24 his relatively lower values of Do for low-molecular-weight dextrans are probably the result of inadequate clarification of his samples. We have also found that clean preparations are absolutely necessary to arrive at a reproducible value for D as measured by P.C.S.The theoretical value of the molecular-weight exponent evaluated by various theoretical treatment^^^ is 0.6, which is reasonable agreement for value for dextran in this study. CONCENTRATION DEPENDENCE ON THE INTRADIFFUSXON COEFFICIENT The concentration dependence of the intradiffusion coefficient, D:, of [3H]T10, [3H]T70 and [3H]FDR7783 is shown as a log-log plot in fig. 4. The intradiffusion coefficient decreases significantly with increasing dextran concentration for all dextran samples studied. An approach to the interpretation of DF is through the treatment of de Gennes.This treatment predicts that above a critical concentration, C*, the flexible polymer chains begins to entangle. Calculation of C* may be made through eqn (5). The values of RG have been obtained through the relationships Do = RTl5.11 x 61q0RGN where qo is the solvent viscosity and Nis Avogadro's number. Values of hydrodynamic radius R, have been obtained from the Stokes-Einstein expression. The calculated values of RG are in good agreement with the measured value of R, for dextran obtained by Granath28 (see table 3). The use of R, to estimate C*, as suggested by1216 DIFFUSION OF DEXTRAN I 0.6 0.8 l m 0 8 0 m o 9 ' 0 8 0. o o 00. 1.0 1 I 0 I I 0 10 100 C/kg rne3 FIG. 4.-A log-log plot of the reduced intradiffusion coefficient D;/(D;)" for tritium-labelled dextrans T10 (a), T70 (0) and FDR7783 (m) against dextran concentration. Adam and Del~anti,~' gives too high a range of C* values.de Gennes predicts that the intradiffusion coefficient, Or, should be independent of concentration up to a value approaching C* and above C* should follow the relationship 0; K C-1.75. This behaviour is commonly handled with a log-log plot as in fig. 4, for which the critical C* is identified with an apparent discontinuation of the variation of 0; with C.Jt is emphasized, however, that such analysis has come under criticism,2s as the nature of the plot lends itself to ambiguous interpretation. In any case, in pursuing the de Gennes treatment we find (contrary to expectations) that 0; values certainly exhibit coilcentration dependence below the critical concentrations evaluated in table 3.In fact, the relative change of 0; with concentration appears maximal at a concentration of 0-30 kg m-3 for T70 and FDR7783. The scaling treatment as applied to the results in fig. 4 is inconclusive as (a) deviations from linearity in the plots appear at relatively high concentrations and (b) the value of the exponent, when an attempt is made to linearise that data, is in the range of 0.7-0.8 for the various dextrans; this value is considerably lower than would be predicted for the movement of a flexible polymer. We also note that the magnitude of the exponent is not sensitively related to the chemical composition and flexibility of molecules under study. Similar values of thePRESTON, COMPER, HUGHES, SNOOK AND VAN MEGEN 1217 TABLE 3.-ESTIMATES OF THE CRITICAL CONCENTRATION, c*, FOR DEXTRANS C*/kg m-3 dextran classification RG/nm R,/nm rangea (RG) rangea (Rh) T10 3.26 2.16 1 18-498 407-1 716 T20 4.83 3.20 7 1-300 256-1030 T70 9.24 6.12 35-146 120-503 FDR7783 14.76 9.78 20-8 1 67-28 1 a Range calculated from eqn (5). exponent have been obtained for the intradiffusion of poly(ethy1ene glycol) and poly(viny1 alcohol) (unpublished) and from a reanalysis of the intradiffusion of albumin.g Indeed the dynamic behaviour of these polymers closely parallels the predicted behaviour of the concentration-dependent movement of spheres and compact particles through networks, where the exponent has a value of 0.75. Experimental values in the range 0.6-0.65 have been obtained for the sedimentation of various spherical particles, including albumin, in poly(ethy1ene oxide) and the transport of albumin in hyaluronate networks30 when reanalysed on the basis of the scaling law.2g There has been one report by Hervet et aL31 which gave the predicted value of the exponent as 1.75 in studies of monodisperse polystyrene in benzene solutions by forced Rayleigh scattering.However, we draw attention to the earlier studies of Park,l* utilizing an albeit more polydisperse polystyrene sample in toluene, where an exponent value of 0.68 was found through measurements of intradiffusion using an open-ended capillary technique (as in this study). There is no clear reason why these two studies should yield such substantially different results. CONCENTRATION DEPENDENCE OF THE MUTUAL DIFFUSION COEFFICIENT COMPARISON OF DIFFUSION COEFFICIENTS OBTAINED BY R.I.M. AND P.C.S.Values obtained by r i m . of the mutual diffusion coefficients of the various dextran fractions, as a function of dextran concentration, are shown in fig. 5. The diffusion coefficients have been measured in solutions covering a concentration range from 1 to 250 kg m-3. For dextrans of molecular weight (uw) > 7 x lo4 a marked concen- tration dependence of D is evident ; with increasing concentration the mutual coeffi- cient becomes greater. On the other hand, for dextrans in the Mw range (1-2) x lo4 the mutual diffusion coefficient is essentially constant. For experiments utilizing similar solutions the diffusion coefficient obtained by P.c.s., as evaulated by the cumulant-fit method, for the various dextran fractions are included in fig.5 (a)-(& There is good agreement between the D values obtained by the P.C.S. and r i m . over the complete concentration range for the T10 and T20 samples, for higher-molecular-weight samples with C < 50 kg m-3, and particularly for FDR7783 at C > 50 kg m-3. In general, the good agreement between the D values obtained using P.C.S. as compared with the boundary relaxation r.i.m., for all four dextran samples measured over a wide range of concentration, immediately establishes the nature of the diffusion coefficient associated with the initial decay of the correlation function as that of a mutual diffusion coefficient. A similar conclusion can be made in the comparison 40 FAR 11218 DIFFUSION OF DEXTRAN l o b O 0 0 0 0 0 0 0 0 t I I I 1 I I I I 1 I 0 (cl 0 0 0 0 0 0 6 6 5 0 0 0.0 50 100 150 200 Clkg m-3 FIG.5.-The mutual diffusion coefficient D of dextran as a function of the mean concentration for dextran T10 (a), T20 (b), T70 (c) and FDR7783 (d): 0, values of the mutual diffusion coefficient obtained by rim.; 0, values of D obtained by P.C.S. as evaluated by second-order cumulant-fitting procedures. of the mutual diffusion coefficient of bovine serum albumin as measured in the ultracentrifuge9 and the diffusion coefficient obtained by quasielastic laser light-~cattering.~~ EFFECT OF POLYDISPERSITY O N THE DIFFUSION COEFFICIENT In that the P.C.S. and boundary relaxation methods yield two different types of average diffusion coefficient, the agreement between the two techniques remains remarkably good.While the prediction of the difference between the two averaged cgefficients is complex and difficult at this stage, the experimentally determined values from the two techniques demonstrate that these differences are not great for the polydisperse samples used in this study. This was confirmed by the P.C.S. analysis of two dextran samples, namely T150 and FDR7783, with approximately the same weight-average molecular weight (1 50000) but with very different values of M,/Mn (1.80 and 1.32, respectively). These results are shown in fig. 6. In this case, values of D for T150 were slightly lower than that obtained for FDR7783, particularly at higher concentrations.PRESTON, COMPER, HUGHES, SNOOK AND VAN MEGEN 1219 v v o 0 0 0 8 0 0 50 100 Clkg m-3 FIG.6.-The effect of polydispersity on the diffusion coefficients, D, obtained by P.C.S. The values of D evaluated by the cumulant-fitting method for dextran T150 and FDR7783 are represented as ('I) and (a), respectively. INFLUENCE OF CONCENTRATION We have provided evidence in this paper that the concentration dependence of the mutual diffusion coefficient of dextran in a good solvent is a function of the molecular weight of the sample. Dextran preparations of T10 and T20 showed effectively no concentration dependence, whereas T70 and FDR7783 exhibited an increase in D with increasing concentration. On the other hand, all the dextran preparations exhibited a marked decrease in the intradiffusion coefficient Dr with increasing concentration.The C* values theoretically predicted and given in table 3 appear too high as applied to the T70 and FDR7783 dextrans. We find no evidence to suggest a 'critical' crossover point between dilute and semi-dilute regions for D measurements outlined in fig. 5. In utilizing the scaling law for mutual diffusion, D cc C", the value of the exponent v is approximately zero for T10 and T20,0.27 for T70 and 0.29 for FDR7783. The values of the exponent are considerably lower than predicted (v = 0.75) for 'cooperative diffusion ' associated with semi-dilute polymer Knowing the virial coefficients of dextran from table 1, we may evaluate the reduced quantity (D/D;t)calc from eqn (4). The variation of the predicted value of (D/Dt)calc is a continuously increasing quantity with increasing dextran concentration.Note that in this evaluation the third virial coefficient dominates the magnitude of the reduced parameter at concentrations > ca. 50 kg m-,, so that the quantity (D/Dt)calc is extremely sensitive to errors in A,, a parameter which is intrinsically difficult to measure. The experimental measurement of the reduced parameter (D/Dt),,p, may be compared with (D/D;t)calc in order to test eqn (4). Owing to the difference in ( D + ) O and Do values outlined in table 2, we have estimated the normalized reduced parameter DID: x (D;t)"/D0 to compare with (D/Dt)calc. For T10 dextran [fig. 7(a)] excellent agreement is obtained between these two reduced parameters up to a concentration of 150 kg m-3.On the other hand, we find for the higher molecular weight dextrans, 40-21220 DIFFUSION OF DEXTRAN 0 50 100 150 0 50 100 Clkg m-3 Clkg m-3 FIG. 7.-Variation of DID;' against dextran concentration for dextran T10 (a), T70 (b) and FDR7783 (c). The solid line is the estimate of (D/Dl+)celc with the use of virial coefficients in table 1 and eqn (4). The use of D from rim. was made in the estimate of DID: x (D;')O/D0 [ =(D/Df),,,J. These values of (D/D:)expt are presented as (0). the parameter D/D: x (D;t>,/Do is consistently less than (D/D;t),,l,. Similar findings were made by ,Laurent el aZ.l0 These authors qualitatively suggested that this discrepancy was the result of the inconsistent use of different averaged parameters describing molecular weight, virialcoefficientsand diffusion coefficients forpolydisperse samples of dextran.However, this observed discrepancy together with conclusions on the scaling treatment of 0;' may also suggest that the value of 0;' is too high for T70 and FDR7783 or that the assumptions involved in the derivation of eqn (4) are not valid for these polymers.PRESTON, COMPER, HUGHES, SNOOK AND VAN MEGEN 1221 This project was supported by the Australian Research Grants Committee (grant nos. D68/16898, D2 73/14137, B78/15168 and DS 79/15252). We acknowledge the expert technical assistance of Gregory Checkley, Wayne Connors and Geoffrey Wilson. W. D. Comper and T. C. Laurent, Physiol. Rev., 1978, 58, 255. B. N. Preston, T. C. Laurent and W. D. Comper, in Glycosaminoglycan Assemblies in the Extracellular Matrix, ed.D. A. Rees and S. Arnott (Humana Press, to be published). J. G. Albright and R. Mills, J. Phys. Chem., 1965, 69, 3120. R. J. Bearman, J. Phys. Chem., 1961,65, 1961. A. G. Ogston, Arch. Biochem. Biophys., 1962, suppl. 1, 39; E. Edmond and A. G. Ogston, Biochem. J., 1968, 109, 569. H. Yamakawa, Modern Theory of Polymer Solutions (Harper and Row, New York, 1971). ' R. Bergman and L. 0. Sundelof, Eur. Polym. J., 1977, 13, 881. L. 0. Sundelof, Ber Bunsenges. Phys. Chem., 1971, 83, 329. R. G. Kitchen, B. N. Preston and J. W. Wells, J. Polym. Sci., Polym. Symp., 1976, 55, 39. lo T. C. Laurent, L. 0. Sundelof, K. 0. Wik and B. Warmegird, Eur. J. Biochem., 1976, 68, 95. M. Daoud, J. P. Cotton, B. Farnoux, G. Jannink, G. Sarma, H. Benoit, R. Duplessix, C. Picot and P. G. de Gennes, Macromolecules, 1975, 8, 804. l2 R. G. Kitchen, Diffiion in Model Connective Tissue Systems (Thesis) (Monash University, 1975). l3 P. G. de Gennes, Macromolecules, 1976,9,587; and review article, P. G. de Gennes, Nature (London), l4 E. Edmond, S. Farquhar, J. R. Dunstone and A. G. Ogston, Biochem. J., 1968, 108, 755; H. Vink, l5 W. E. Trevelyan and J. S. Harrison, Biochem. J., 1952,50, 298. l6 J. M. Creeth, J. Am. Chem. SOC., 1955, 77, 6428. L. J. Gosting and H. Fujita, J. Am. Chem. Soc., 1957, 79, 1359. G. S . Park, Symp. Macromolecules, IUPAC Meeting, Weisbaden, 1959, vol. 11, paper AS. 1979, 282, 367. Eur. Polym. J., 1971, 7, 1411. lD J. S. Anderson and K. Saddington, J. Chem. Soc., 1949, S381. 2o L. Nanis, M. Litt and J. Chen, J. Electrochem. Soc., 1973, 120, 509. 21 R. J. Pecora, J. Chem. Phys., 1968, 49, 1036. 22 M. B. Weissman, J. Chem. Phys., 1980, 72, 231. 23 P. N. Pusey, Milan ConJ Light Scattering Fluids of Macromolecular Solutions, ed. V. Giogo, M. Corti and N. Cigilio (Plenum Press, New York, 1980); K. Gaylor, I. Snook and W. van Megen, J. Chem. Phys., 1981, 75, 1682. 24 D. B. Sellen, Polymer, 1975, 16, 561. 25 M. Daoud and G. Jannink, J. Phys. (Paris), 1976, 37, 293; S. F. Edwards, Proc. Phys. SOC., 1965, 28 K. Granath, J. Colloid Sci., 1958, 13, 308. 27 M. Adam and M. Delsanti, Macromolecules, 1977, 10, 1229. 28 G. D. Patterson, J. P. Jarny and C. P. Lindsay, Macromolecules, 1980, 13, 668. 29 D. Langevin and F. Rondelez, Polymer, 1978, 19, 875. 30 T. C. Laurent, and A. Pietruszkiewicz, Biochim. Biophys. Acta, 1961, 49, 258. 31 H. Hervet, L. Leger and F. Rondelez, Phys. Rev. Lett., 1979, 42, 1681. 32 B. D. Fair and A. M. Jamieson, J. Colloid Interface Sci., 1980, 73, 130. 85, 613. (PAPER 1/859)

ISSN:0300-9599    

DOI:10.1039/F19827801209

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1982,78, 1223-1236 Solvent Effects in the Electron Spin Resonance Spectra of Semiquinones B Y DOLORES M. HOLTON~ AND DAVID MURPHY* Department of Chemistry, Bedford College, Regent's Park, London NW 1 4NS Received 1st June, 1981 Solvent effects on the e.s.r. spectra of semiquinones in mixtures of H,O + hexamethylphosphoramide (HMPA), H,O +dimethylsulphoxide (DMSO), H20 + dimethylformamide (DMF), H,O + EtOH and EtOH + HMPA are reported and related to solvent basicity and radical structure. In many instances quantitative studies were possible and it is demonstrated that true thermodynamic equilibria are being studied. The significance of the measured equilibrium constants is discussed and an interpretation in terms of preferential solvation by the aprotic solvent is rejected in favour of one in which solvent-solvent and solvent-radical interactions operate simultaneously.In H20 + HMPA mixtures, for instance, competition between the radical and HMPA for water molecules is envisaged, the overall result being measured by e.s.r. This suggestion is compatible with the observed linear dependence of the coupling constants on [HMPA]/[H,O] and on measures of solvent polarity such as dielectric constant and the Kosower Z value. In a previous publication' it has been shown that a simple theoretical model due to Gendell et aL2 is applicable to the equilibrium between the two forms of a semiquinone present in a mixed solvent system: S, + RS, RS, + S, (1) where RS, is the radical solvated by solvent S, and RS, is the radical solvated by solvent S,.The equilibrium constant ( K ) for eqn (1) is given by Exchange between the two forms was always fast in the systems studied, so that K can be expressed in terms of a coupling constant, a, where a has the values a, in solvent S, and a, in solvent S,: Values for K are obtained either from sijpoid plots of a against log([S,]/[S,]) or, in more favourable cases, from one of the linear plots Ia-a,l-l against [S,]/[S,] and la- a,/-, against fS,]/[S,]. In the present paper the range of semiquinones for which thermodynamic data can be obtained is extended and the method is applied to systems other than S, = H,O, S, = HMPA. t Present address: University Chemical Laboratory, Lensfield Road, Cambridge CB2 1 EW. 1223TABLE 1 .-E.s.R.PARAMETERS (a f 0.031 10- T) AND EQUILIBRIUM CONSTANTS FOR SEMIQUINONES IN THE H20 + HMPA SYSTEM (i) substituents in 0- 4 ,"o; P ? w I m m 0- solvent a, a3 a5 '6 K P parameter used cd n I I I I I I 2-CH3 I 2-OCH3 1 2,5-But2 I A 2,6-But, I :20 :,0 [ZO i,O i20 :MPA [MPA [MPA [MPA [MPA B 2,6-(OCH3), HZ0 HMPA 2.38 2.38 2.45 2.45 aCHs = 2.10 1.73 a C H s = 1.77 2.05 aOCHs = 0.82 0.52 aOCHs = 0.47 1.16 2.08 2.33 1.31 2.25 aoCHs = 0.80 1.47 aOCHs = 0.50 1.92 aCHa = 2.10 1.87 aCHs = 1.85 2.22 2.38 2.45 2.60 2.73 3.60 3.48 1.31 2.25 1.47 1.92 1.87 2.22 2-38) 1.40b 2.45 2*39} 0.3Y 0.929 2.43 0-67c Om9"32 1.94 2.31 1 { 0.68d 0.993 2-08 } 0.07b 2.33 { 0.8W 1.16c 0-999} 0.996 aoCHs = 0.80 I 0-58c 0.68d 0.975} 0.995 0*986 } aOCHs = 0.50 0.38' '1 O.ad 0.980(ii) substituents in 0- Q5 4 K ra parameter used C H2O D 3-CH3 H2O E 3-OCH3 HZO F 4-CH3 H2O G 4,5-(OCH3), H2O H 4,5-OCH20 H2O HMPA HMPA HMPA HMPA HMPA HMPA HMPA 3,4,6-(OCH3), H20 0.75 1.56 aCH3 = 0.63 aoCH3 = 0.65 aOCH3 = 0.47 0.07 1 .oo - 0.32 0.33 -0.33 0.50 uCHs = 1.06 aoCH3 = 0.00 a o C H 3 = 0.00 3.69 3.34 2.85 2.83 1.25 1.90 acH3 = 4.89 uCHs = 3.85 uOCH3 = 1-05 aoCH3 = 0.85 ( I C H 1 = 4.22 uCHa = 2.39 aoCH3 = 0.90 u o C H 3 = 0.50 3.69 3.34 4.13 3.54 4.76 3.78 3.89 3.55 aOCH3 = 1.05 aOCH3 = 0.85 a C H 2 = 4.22 4CH2 = 2.39 1.15 2.13 0.75 1.56 0.28 1.06 0.46 0.98 } { K3:: 1.69 } { ;:z } { ::::: - 0.59 -0-32} 0.33 0.80" 0.22" 0.23d Oe50 ' 0.30" -0.33 0.998 0.997 1 0.989 0-990 1 0.994 0-994 1 0.993 0.983 0.990 0.996 0.972 0.995 a, + Q, +a, Q5 a 3 aCH3 2Q3,6 aCHz 2a3.5 a3 +a, + a, + a, p parameter used solvent a29 a 3 a57 a 8 as7 Q7 K 3.20 0.66 0.66 0.13" 0.965 2a,,, @ 0: H20 HMPA 3.33 0.24 0.65 } 0- a Correlation coefficient from linear regression analysis. Equilibrium constants from plots: Q against log ([S,]/[S,]>; la-ull-l against [Sll/[S21; 10- against ks2]/s~].1226 0.5- 50.0- 2 -2 I -0.5- E.S.R. SPECTRA OF SEMIQUINONES EXPERIMENTAL MATERIALS 1,4-dihydro~y-2,6-dimethoxybenzene,~ 1,2-dihydroxy-4,5-methylenedi~xybenzene~ and 1,2- dihydroxy-3,4,6-trimethoxybenzenes were prepared as in the cited literature.Other precursors and solvents were commerical materials. All compounds used were purified by the usual methods and had physical constants which agreed well with those of the literature. E.S.R. SPECTRA Mixtures of the two solvents were made up by weight fraction.The radicals were generated in the usual way by autoxidation,s using KOBut (solid) as the base. E.s.r. spectra were run, using a static system, on a Varian E4 spectrometer. RESULTS THE H,O+HMPA SYSTEM Table 1 contains equilibrium constants obtained in H,O + HMPA. The straight lines which best fitted the data were found by linear regression, the correlation coefficient, r, giving the quality of the ‘fit’. Where no value of r is given the magnitude of the changes incurred on altering the solvent was too small to allow meaningful equilibrium constants to be obtained from the linear plots; instead an approximate value was determined from the appropriate sigmoid curve. Wherever possible the parameter undergoing the largest change from water to HMPA was employed in calculating K.For methyl- and methoxy-substituted radicals, significant differences in the values of K were found if a splitting due to an alkyl proton rather than a ring proton was used in the calculation. Compare, for example, values given in table 1 for radicals (B) or (F) when different coupling constants are employed. In constrast, the total width and the parameter 2u3,, for 1,2-benzoserniquinone (C) yield the same value, as do uCHp and 2u3,6 for 4,5-methylenedioxy- 1,2-benzosemiquinone (H). At present no explanation for this behaviour may be advanced. The variation of a coupling constant over the entire range from pure water to pure HMPA reveals certain features which might have been overlooked had only the extreme values been available.It is known, for instance, that a, in radical (E) is negative when the radical is generated in water; this is confirmed by SCF calculations and I 1 I I I I -2.0 -1.0 0.0 1.0 2 .o 3.0 log (IH201 /[HMPAl) FIG. 1 .-Variation in a, for 3-methoxy-1 ,Zbenzosemiquinone (E) with solvent composition.D. M. HOLTON AND D. MURPHY 1227 - 2 x 1 0 - 4 ~ fd) FIG. 2.-Spectra of 4,5-methylenedioxy-l ,Zbenzosemiquinone (H) in H,O + HMPA. (a) Pure water, u ~ , ~ = (-)0.33 x lo-’ T; (b) [H,O]/[HMPA] = 19.33, u ~ , ~ = (-) 0.17 x lo-’ T; (c) [H,O]/[HMPA] = 6.51, a3,@ = 0.0; (d) pure HMPA, a3,,, = 0.50 x lo-* T. empirical graphical procedures. By observing the smooth change in a, with solvent composition (fig. 1) it is apparent that it changes sign. The same phenomenon is observed for radicals (G) and (H); some spectra for (H) are presented in fig.2. Note that five of the seven ‘best’ radicals in table 1, as suggested by the values of t, are o-semiquinones [radicals (C)-(F), (H)], the remaining two being 2,6-di-t-butyl- and 2-methoxy-p-benzosemiquinone. Omitting the last radical, in which the changes in coupling constants are too small in other solvents, these six radicals were selected for further study. EQUILIBRIUM CONSTANTS IN OTHER SOLVENT SYSTEMS Having verified that it leads to reasonably consistent equilibrium constants in H20+HMPA (table l), the treatment was extended to H20+DMS0, H,O+DMF, EtOH+HMPA and EtOH+H,O. Initially, a change in the aprotic solvent was studied : tables 2 and 3 contain equilibrium constants determined in the H20 + DMSO and H20 + DMF systems, respectively. In general, the linear regression [with the exception of radical (C)] is less satisfactory than in H20 + HMPA, but is nonetheless acceptable.Table 4 lists equilibrium constants obtained in the EtOH + HMPA system. Serious practical difficulties were encountered in this system due to the formation of secondary and dimeric radical species; these produced spurious lines in the e.s.r. spectra of the required radicals towards the ethanol end of the solvent range. Owing to the smaller variations in the coupling constants in EtOH + H20 compared with protic + aprotic mixtures, coupled with the difficulties inherent in obtaining ‘clean’ spectra in ethanol, it was not possible to construct linear plots for the majority of radicals in table 5.L h) h) 00 TABLE 2.-EQUILIBRIUM CONSTANTS IN H2O + DMSO (i) radical solvent a3 a4 a5 a, K r parameter used 0- 0- 1.31 2.01 H2O DMSO (ii) substituents in 0- &O- solvent a3 a4 a5 a, K r parameter used C D 3-CH3 E 3-OCH3 F 4-CH3 2a3,s 0.75 3.69 3.69 0.75 0.57“ 1.39 3.45 3.45 1.39 1 {0.57d 1000 DMSO aCH3 = 0.63 2.85 4.13 0.28} { C):M;i DMSO aCH3 = 1.00 2.80 3.70 1.00 0.993 DMSO aoCHs = 0.60 1.60 4.00 0.18 DMSO H@ H2O H2O H2O 0.997 } total width aoCH3 = 0.65 1.25 4.76 -0.59) 0.63d 0.992 a, 0.07 ~ C H ~ = 4.89 3.89 0.98 0.24“ 0.956) total width 0.80 aCH3 = 4.08 3.62 1.51 1 {0.26d 0.954 U d U 2: 0 a-d As in table 1.TABLE 3.-EQUILIBFUUM CONSTANTS IN H 2 0 + DMF ~ (i) radical solvent a3 a4 a5 ' 6 K r parameter used 1.31 2.12 2.12 le31 ) I 0- P (ii) substituents in ic: X 0 r c3 0 2, * 21 '6 K r parameter used solvent a3 a4 a5 0.75 3.69 3.69 0.75 0.54c 0.991) 2u3,6 U 3.48 1.28 { 0.45d 0.999 P C H2O DMF 1.28 3.48 DMF uCH3 = 1.06 2.90 3.66 0.98 } { 0.48d 0.998 z d P X D 3-CH, H20 uCH3 = 0.63 2.85 4.13 0.28 0.58c 0*998} total width 1 .07c cd E 3-OCH3 H2O DMF uoCH3 uOCH3 = 0.65 0.60 1.25 1.75 4.76 4.00 -(I:):} [0.91d 0.72c (I:;:} 4 " 0.41d ::;;:} H2O 3.89 0.98} 0.97c 0.978 a3 0.07 uCH3 = 4.89 3.65 1.60 F 4-CH.3 DMF 0.89 uCH3 = 4.01 -0.33 uCHz = 4.22 uCHe = 4.22 -0.33) { 0:71: uCH~ 0.43 uCHZ = 2.52 uCHz = 2.52 0.43 178 2u3,6 H 4,5-OCHzO H2O DMF a-d As in table 1.TABLE ~.---~QUILIBRIUM CONSTANTS IN EtOH -t HMPA (i) radical solvent a3 a4 a5 a6 K T parameter used 0- EtOH 1.21 2.25 0- ~~ (ii) substitutents in 0- a4 a5 a6 K r parameter used C EtOH D 3-CH3 EtOH HMPA HMPA F 4-CH, EtOH H 4,5-OCH,O EtOH HMPA HMPA 0.93 3.59 3.59 0.932 { WIi 1.56 3.34 3.34 1.56 aCHs = 0.80 2.80 4.07 0.44 0.87' uCHs = 1.06 2.83 3.54 1.06) (0.53" 0.59' aoCHJ = 0.65 1.30 4.59 -::I} 10.49" aoCHs = 0.47 1.90 3.78 0.67' 0.65" 0.33 aCHs = 4.47 3.70 1.10 0.78' 1.00 aCHs = 3.85 3.55 1.69 1 { 0.72" -0.08 aCHl = 3.64 3.64 -0.08 0.65' 0.50 aCHI = 2.39 2.39 0.50) (0.66" m m a-d As in table 1.TABLE 5.-EQUILIBRIUM CONSTANTS IN EtOH 4- H,O (i) radical solvent a3 a4 '6 '6 K r parameter used 1.31 1.21 le31 1.21 } 0.2b 2a3.5 0- (ii) substituents in 0- a4 a5 K r parameter used C H2O 0.75 3.69 3.69 0.36c 0.994) 2a3,6 EtOH 0.93 3.59 3.59 "0;: 1 \0.34d 0.998 D 3-CH3 H2O E 3-OCH3 H@ F 4-CH3 H2O H 4,5-OCH20 H2O EtOH EtOH EtOH EtOH aCHs = 0.63 aCHa = 0.80 aoCHs = 0.65 aoCHs = 0.65 0.07 0.33 - 0.33 - 0.08 2.85 4.13 2.80 4.07 1.25 4.76 1.30 4.59 aCHs = 4.89 3.89 aCHs = 4.47 3.70 aCHr = 4.22 aCHI = 4.22 aCHI = 3.64 a C H t = 3.64 0.44 - 0.40 1.10 - 0.08 0.45b 0.54b 0.30b 0.2Y { 0.21d 0.996 0.999 total width total width total width 2aCHo a-d As in table 1.1232 E.S.R. SPECTRA OF SEMIQUINONES TABLE 6.-CORRELATION OF COUPLING CONSTANTS OF RADICALS (A), (c) AND (F) WITH SOME SOLVENT POLARITY PARAMETERSa solvent 1. HMPA 29.6 2. acetone 20.70 3. DMF 37.0 4. DMSO 46.48 5. sulpholane, 43.3 TMSO, 6. ethanol 24.55 7. methanol 32.7 8. fomamide 11 1.0 9. water 78.39 10. 2-methyl-propan-2-01, 12.47 1 1. ethanediol 37.7 12. propan-2-01 19.92 13.butan-1-01 17.5 1 14. propan-1-01 20.33 t-BuOH 62.8 65.5 68.4 71.1 77.5 79.6 83.6 83.3 94.6 71.3 75.1 76.3 77.7 78.3 40.9 42.2 43.8 45.0 44.0 51.9 55.5 56.6 63.1 43.9 56.3 48.6 50.2 50.7 4.50 3.75 4.24 4.02 3.60 2.42 2.35 2.95 2.62 2.36 2.35 3.58 2.40 2.48 3.12 2.56 2.78 2.48 - 1.88 1.84 1.96 1.50 1.80 1.77 1.91 - - 1 .oo 0.89 0.80 0.60 - 0.33 0.3 1 0.38 0.07 a E , Z/kcal mol-l and ET(30)/kcal rno1-l are taken from ref. (7). Values of 2a3, &, 2a,, and a3 in T. CORRELATION OF COUPLING CONSTANTS WITH SOME MEASURES OF SOLVENT POLARITY In table 6 the hyperfine splitting constants of three representative semiquinones (A), (C) and (F) are presented, together with some physical properties’ of the solvents in which the spectra were determined. Overall, the correlation between the coupling constants and E , the dielectric constant, is poor; this has been found for other radicals by previous workers.However, for the solvents H,O, DMF, DMSO and HMPA, which are of particular interest in the present study, the correlation between a and E is 0.995,0.951 and 0.990 for radicals (A), (C) and (F), respectively. The success of this correlation suggests that the coupling constants of a given radical in any of these solvent systems are a linear function of the dielectric constant. This was tested on data for radical (C) in H20 + EtOH and H20 + DMSO mixtures. Over the entire concentration range from water to ethanol, there is a straight-line relationship between 2a,,, and E (r = 0.998). This is also true of H,O + DMSO solutions from pure DMSO to [DMSO]/[H,O] = 2.0 ( r = 0.993).When the Kosower 2 value8 is used as the measure of solvent polarity, the correlation is generally good and is better than that found with E . Values of r for radicals (A), (C) and (F), respectively, in the solvents water, DMF, DMSO and HMPA are 0.999, 0.976 and 0.998. For radical (A) that between the Dimroth-Reichardt polarity parameter E , (30)9 and a is 0.998. The excellent correlation suggests that the e.s.r. studies are measuring similar effects to those measured by the 2 value of the solvent. The effect of changing solvent polarity within the series of mixtures used for a ‘run’ was investigated for radical (C). From pure ethanol to [H,O]/[EtOH] = 2.0, the hyperfine splitting is a linear function of 2 (r = 0.989), while in H,O+DMSO theD.M. HOLTON AND D. MURPHY 1233 correlation ( r = 0.999) is from pure water to pure DMSO. 2 values are not currently available for other protic + aprotic mixtures. Correlations between e.s.r. parameters and 2 values have been noted for other radicals in pure solvents but their success appears to depend on the nature of the radical and the type of solvent studied. For p-benzosemiquinone,lO for example, a('") and a(170) are linearly related to the 2 values of solvents such as water, ethanol and DMSO, whereas only a poor correlation is observed for 2,6-dimethyl-p-ben~osemiquinone.~~~ l2 Here the major additional factor isconsidered to be steric. In the present study, aH for the 2,6-di-t-butyl-substituted radical (A), although a smooth function of 2, is not linearly related to it.It is likely that the steric explanation applies in this case also. DISCUSSION RADICAL-SOLVENT INTERACTIONS In the course of this investigation, the solvent dependence of radicals in the solvents water, ethanol, DMSO, DMF and HMPA has been studied qualitatively and, in favourable cases, quantitatively. In the aprotic solvent HMPA differences from the water hyperfine splitting constants are at a maximum for all radicals. Water is a good anion solvator by virture of its hydrogen-bonding ability, whereas in HMPA anion solvation is mainly via ion-dipole and ion-induced-dipole forces. Due to steric hindrance around the phosphorus atom anions are practically unsolvated in HMPA. With the other aprotic solvents, changes from the water values are smaller, presumably as interactions can occur to some extent between the anion and DMSO or DMF.The results demonstrate that the anions are only weakly solvated in aprotic solvents, but the protic solvents employed form hydrogen bonds to the basic sites of sufficient strength to perturb the semiquinones in a readily detectible way. l3 These perturbations are manifested as changes in coupling constants with solvent. APPLICATION OF THE GENDELL, FREED AND FRAENKEL MODEL TO SEMIQUINONES I N MIXED SOLVENTS Tables 1-5 demonstrate conclusively that the simple model2 applies to the equilibrium involving semiquinones in mixed solvents [eqn (l)]. Data from all five solvent systems investigated comply with this equation and graphs are of the form predicted by the theory for each case investigated.The consistency of the model is supported by the good agreement between equilibrium constants obtained using different plots: a against log([S2]/[Sl]), la - q1-l against [S,]/[S,] and la - a21-l against [S,]/[S]. Therefore, a single equilibrium is being measured throughout the range. Table 1 shows that the method is sensitive enough to differentiate between differently substituted semiquinones. For example, for radical (C), K has been obtained from two different linear plots and the values differ by 0.06, which is within experimental error. The difference between K for radical (C) in H20+HMPA and that for any of the other radicals is, in general, greater than the experimental error and the values of K are reproducible. In the case of radicals (G) and (H), the equilibrium constants are distinct although the substituents have similar effects on the spin density of a radical.Different 2,6-sulrstituents also lead to separate equilibrium constants (see table 1). MAGNITUDE OF THE EQUILIBRIUM CONSTANTS Tables 1-5 show that, with the particular exception of 2,6-disubstituted-p- semiquinones, which are asymmetrically solvated, equilibrium constants are generally less than unity. Comparing values obtained using the same parameters, K for1234 E.S.R. SPECTRA OF SEMIQUINONES H,O + HMPA is usually quite different from that for other aqueous + aprotic solvent systems. This is particularly marked for radicals (A) and (C), where equilibrium constants for H,O+DMF/DMSO are similar. In DMF and DMSO, of course, the radicals are not 'free' as they are in HMPA.The observed order of equilibrium constants changes with the nature of the radical : (A), (E), (F) K(H,O + HMPA) < K(H,O + DMSO) < K(H,O + DMF) K(H,O+HMPA) < K(H,O+DMSO) x K(H,O+DMF) K(H,O + HMPA) x K(H,O + DMSO) < K(H,O + DMF) (C) (D) K is, therefore, a function of the substitution pattern as well as the solvent combination. The immediate implications of equilibrium constants less than unity are rather surprising. In the original model,, K = 10 defines a situation in which the complex with solvent S, is considerably stronger than that with S, [eqn (2)]. A value of unity corresponds to unselective solvation, such as has been found for nitrobenzene anions.14 For the H,O + HMPA system, for example, K < 1 could be taken [eqn (4)] R..*HMPA+H,O~R"-H,O+HMPA K < 1 (4) as preferential solvation of the semiquinone radical by HMPA, but in view of the previous discussion this is an unacceptable conclusion.Moreover, values in H,O + HMPA are lower than those in H,O + DMF/DMSO, implying that solvation by HMPA is greater than that by DMSO or DMF, which is again unacceptable. Also, results from EtOH+H,O suggest [eqn (S)] that hydrogen bonding to the radical is stronger for ethanol than water: R. * .EtOH +H,O R. * .H,O + EtOH K < 1. ( 5 ) This is inferred, also, when values for H,O+HMPA and EtOH+HMPA are compared as K(Et0H + HMPA) is greater than K(H,O + HMPA). Equilibrium constants less than unity lead to positive AG* values, i.e. the reaction is unspontaneous. It is likely, therefore, that some other effect is operative in this system additional to that represented by eqn (1).SOLVENT-SOLVENT INTERACTIONS The results from the trihydroxybenzenes and n.m.r. studies are strong evidence that the additional effect overlooked in eqn (1) is an interaction between the two solvents. E.s.r. results from trihydroxybenzenes in H,O + HMPA showed the solvent structure to change in a definite manner, so that radical dianions cannot persist at high HMPA concentrations.l N.m.r. studies1*15 confirmed that this change was due to the formation of the relatively stable complex HMPA * 2H,O. Similarly, strong solvent- solvent interactions and complex formation were found for H,O + DMF, H,O + DMSO, EtOH + DMF, EtOH + DMSO and EtOH + HMPA.This would explain why the equilibrium constants found by e.s.r. are positive: there is actually competition for water molecules between the radical and HMPA and it is the result of this competition which is being measured by e.s.r. This conclusion is supported by studies on other types of radical. For the reaction : HMPA. * *H-OH+PhNOi-+HMPA+PhNOi-. * -H-OH K = 0.9, indicating that the concentration of hydrogen-bonded HMPA is greater than that for the hydrogen-bonded ion,lg an expected result as HMPA forms relatively strong hydrogen bonds.,' Equilibrium constants less than unity for t-butylnitroxide in mixed aqueous solvents1* were interpreted in terms of preferential solvation byD. M. HOLTON AND D. MURPHY 1235 organic solvents such as dioxan and alcohols.However, using di-t-butylnitroxide as a probelS this concept was shown to be totally misleading. Interaction with aprotic solvents is dipolar in nature. As the concentration of this type of solvent is increased, hydrogen bonds to the radical are lost due to the strong affinity of the basic solvent for hydrogen bonds, rather than to increased radical-aprotic-solvent interactions. Thus both radical-solvent and solvent-solvent interactions must be considered. The reliance of a on [H,O]/[HMPA] might appear to suggest that the equilibrium between the two radical forms is effectively independent of any interaction between the solvents, i.e. that eqn (1) stands. In the ideal case of zero solvent-solvent interactions such a relationship is expected. However, the positive AG* values deny this. Rather, there appear to be two simple situations for a semiquinone: either it is hydrogen bonded or it is not.The effects on spin density distribution within the radical would be expected to depend most critically on the strength of the primary hydrogen bond, that to the radical; thus it would be insignificant whether the radical were hydrogen bonded to a water molecule or to the HMPA * 2H20 complex, and la - all-1 could still be a linear function of [H,O]/[HMPA]. The structure of HMPA-2H20 is probably the same as that suggested for H20 + DMSO interactions:20 .* The first water molecule is bonded to the oxygen atom of the aprotic solvent (lo solvation), whereas the second water molecule is attached to the first by secondary solvation (2O), as shown above.Although it is often assumed that 1' solvation is complete before the onset of 2 O solvation, i.r. studies on tetra-alkylammonium halides21 have established that this is incorrect. The structure proposed for HMPA-2H20 has the capacity to form further hydrogen bonds. Depending on the relative concentrations, these can be either to the radical or to other water molecules. Thus the measured coupling constants are averages resulting from exchange between hydrogen-bonded radicals, either bonded to water or to HMPA - 2H20, and radicals solvated through dipolar interactions with HMPA. However, further details are not obtainable from the present results. This explanation is compatible with the finding that the coupling constants correlate simultaneously with E and 2, one a bulk, macroscopic property of the solvent and the second a measure of interactions at the molecular level.If the present picture of hydrogen bonding to the radical by HMPA-2H20 or water itself is accepted, the observed correlations are consistent : a must depend on the nature of the lo bond, i.e. on what is happening at the radical and therefore on 2, and also on the interactions throughout the solvent as a whole and therefore on E. Whatever the physical significance of the equilibrium constants it is clear from the linear relationships that unique values are being measured. Thermodynamic consid- erations confirm this. THERMODYNAMIC CONSIDERATIONS If the measured K values are true thermodynamic equilibrium constants, then AG* values calculated from them must be additive.Therefore, if two equilibrium constants are known, a third can be calculated and compared with the experimentally observed value. Constants for radicals in H 2 0 + EtOH were obtained in this manner (table 7).1236 E.S.R. SPECTRA OF SEMIQUINONES TABLE 7.-hEDICTED EQUILIBRIUM CONSTANTS FOR RADICALS IN HZO 4- EtOH K(H,O + HMPA) K(Et0H + HMPA) K(Et0H + H,O) K(Et0H + H,O) radical experimental experimental experimental calculated source of K experimental (C) 0.29 0.84 0.35 0.36 h,,, against log{[SJ/[S,]} (D) 0.19 0.71 0.45" 0.27 a, against log{[S,]/[S,]} (E) 0.40 0.60 0.54" 0.67 a6 against log{[SJ/[S,]} (F) 0.13 0.56 0.30 0.23 total width against log{[S,]/[S,]} (H) 0.16 0.63 0.25 0.22 aCHI against log {[S,]/[S,]) " K(H,O + EtOH) from total spectral width against log{[S,]/[S,]] plot. The limiting factor to accuracy in these calculations is the H20 + EtOH experiment. Values in this system are probably only correct to f0.2 due to the small changes involved and the fact that K can only be determined from the less reliable a against log ( [ S , ] / [ S , ] } plot; for other systems K is estimated to be correct to f 0.1.For con- sistency, only values obtained from log plots have been'used and, where possible, the same parameter has been employed in the calculation of K. In all cases the agreement between predicted and experimental values of K is good and, in fact, disparities are only greater than 0.1 when values obtained from different plots have had to be compared [radicals (D) and (E)].This supports the suggestion that a single, true equilibrium is being observed in the e.s.r. studies, that between hydrogen-bonded and non-hydrogen- bonded radicals. D. M. H. thanks the S.R.C. for a research grant. W. T. Dixon, D. M. Holton and D. Murphy, J. Chem. SOC., Faraday Trans. 2, 1978, 74, 521. E. Chapman, A. J. Perkin and R. Robinson, J. Chem. Soc., 1927, 3015. F. Dallacker, W. Edelman and A. Weiner, Annalen, 1968, 719, 112. W. Baker, J. Chem. SOC., 1941, 662. W. T. Dixon, P. M. Kok and D. Murphy, Tetrahedron Lett., 1976, 623. C . Reichardt, Solvent Eflect in Organic Chemistry (Verlag Chemie, Weinheim, 1979). E. M. Kosower, J. Am. Chem. SOC., 1958,80,3253; J. Chim. Phys., 1964,61,230; E. M. Kosower and M. Mohammad, J. Am. Chem. SOC., 1968,90, 3271 ; 1971,93,2713; J. Phys. Chem., 1970,74, 1153. K. Dimroth, C. Reichardt, T. Siepmann and F. Bohlman, Annalen, 1963, 661, 1 ; C. Reichardt, Annalen, 1971, 752, 64. * J. Gendell, J. H. Freed and G. K. Fraenkel, J. Chem. Phys., 1962,37, 2832. lo T. E. Gough and M. C. R. Symons, Trans. Faraday SOC., 1966, 62, 269. l1 T. A. Claxton, J. Oakes and M. C. R. Symons, Trans. Faraday SOC., 1967, 63, 2125. l 3 M. C. R. Symons, Pure Appl. Chem., 1977, 49, 13. l4 J. M. Gross and M. C. R. Symons, Trans. Faraday SOC., 1967, 63, 21 17. l5 D. M. Holton and D. Murphy, unpublished results. G. R. Stevenson and H. Hidalgo, J. Phys. Chem., 1973, 77, 1027. l7 T. Olsen, Acta Chem. Scand., 1970, 24, 3081. l8 G. Stout and J. B. F. N. Engberts, J. Org. Chem., 1974, 39, 3800. l9 Y. Y. Lim, E. A. Smith and M. C. R. Symons, J. Chem. SOC., Faraday Trans. I , 1976, 72, 2876. *O J. M. Harvey, M. C. R. Symons and R. J. Naftalin, Nature (London), 1976, 261,435; J. M. Harvey, S. E. Jackson and M. C. R. Symons, Chem. Phys. Lett., 1976,37,551; J . Chem. SOC., Faraday Trans. I , 1980, 76, 256. P. S. Gill and T. E. Gough, Trans. Faraday SOC., 1968, 64, 1997. *l M. C. R. Symons and S. E. Jackson, J. Chem. Soc., Faraday Trans. I , 1979, 75, 1919. (PAPER 1/881)

ISSN:0300-9599    

DOI:10.1039/F19827801223

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1, 1982, 78, 1237-1247 Effects of Pressure and Surface Initiation Efficiency on the Flowing H,O, + NO, + CO + N, Chain Reaction System BY GARY J. AUDLEY, DONALD L. BAULCH AND IAN M. CAMPBELL* School of Chemistry, The University, Leeds LS2 9JT Received 11 th June, 1981 The yields of CO, from the H,O, +NO, + CO reaction system in flowing camer gas at 298 K have been measured at total pressures ( P ) of 13.3,40.0 and 93.3 kPa. The variations of the yield as functions of [NO,] and [S] (S = acetaldehyde or diethyl ether) have been shown to indicate that the efficiency of conversion of H,O, to OH radicals in the heterogeneous initial step H,O, + NO, 4 OH + HNO, (1) is independent of total pressure for a particular surface activity. The results are consistent with the pressure dependence of k , for the propagation step CO+DH 4 CO,+H (2) (4) indicated by other work.For the rate constant of the termination reaction OH +NO,( + M) -+ HNO,( + M) low apparent values (k;) are indicated compared with literature values of k,. At P = 40.0 kPa, for three surfaces of apparent efficiencies of ca. 20%, ca. 53% and ca. 81 % in converting H,O, to OH in reaction (l), k;/k, values were ca. 0.45, ca. 0.65 and cu. 0.85, respectively. This trend is interpreted in terms of decreasing activity of these surfaces for adsorption of H,O,, producing decreasingly localized reaction zones : depletion of [NOJ in the reaction zone depresses k i / k , while increased surface concentrations of H,O, decrease the apparent H,O, to OH conversion efficiency by inducing an additional route for destruction of the active centre involved in reaction (1).Hydroxyl (OH) radical reaction kinetics are of outstanding interest at present due to recognition of the major roles of OH in atmospheric and combustion chemistry. Also, as pointed out in a recent review,l an upsurge in laboratory studies of OH kinetics in the past decade has resulted from developments in methods of generating and detecting the radical. Of newer methods of producing OH radicals, the thermal source in H,O, + NO, + CO gas-phase mixtures at ambient temperatures, developed in this laboratory,2$3 has shown considerable potential in the measurement of rate constants for reactions of OH with vapour-phase molecules such as alcohols,* e s t e r ~ , ~ aldehydess and nitro-compo~nds,~~ * many of which were first measurements.In all of our previous work the conditions used were a partial pressure of hydrogen peroxide considerably less than the saturation vapour pressure at ambient temperature, an excess of nitrogen dioxide and a large excess of CO or a CO+N, mixture, at a total pressure of d 13.3 kPa. The initiation step, occurring on the wall of the reaction vessel, was proposed to be (1) It was shown3 that coating of the walls with boric acid allowed achievement of a ca. 100% efficiency for conversion of H,O, into,OH in step (l), although usually the efficiency was lower. In this work we set out to extend the total pressure range up to approaching H,O, +NO, -, OH + HNO,. 12371238 H,O,+NO,+CO+N, CHAIN REACTION SYSTEM atmospheric. In doing so, one of our interests was in whether the stoichiometry of the heterogeneous process represented by reaction (1) was independent of total pressure or if some other reaction pathway might be induced.This is of considerable importance if reaction (1) is to be applied to the study of OH-induced reaction kinetics under conditions representative of the troposphere at large. It is also of importance to our understanding of the reaction system to attempt to gain some understanding of the reasons why the surfaces often induce H20, to OH conversion efficiencies of less than 100%. We have encountered three surfaces of different efficiencies in this work. Subsequent to initiation by reaction (I), the propagation steps of the chain cycle with OH as the chain carrier are CO+OH +CO,+I-I H+NO,+OH+NO while with excess NO, present the overwhelming termination step is postulated to be (4) The second-order rate constants k, and k, are likely to be pressure-dependent across our total pressure range of 13-94 kPa in CO + (N,) carriers on the basis of a substantial number of studies reported in the literature.' Accordingly, another aim in studying the H,O, +NO, + CO + (N,) system over this pressure range was to examine to what extent the measured yields of CO, from reaction (2) were consistent with the literature values of k, and k, for each pressure.OH + NO,( + M) + HNO,( + M). EXPERIMENTAL In this work we used the flowtube system which has been described in detail before,8 so that only the essential aspects are given here.The main flowtube was constructed of Pyrex glass, with an internal diameter of 24 mm, and was 1.15 m long. It was contained in a jacket through which water, thermostatted at ca. 298 K, was circulated continuously. An inner liner of Pyrex glass, of internal diameter 21 mm and of length 1.0 m, fitted flush within the main flowtube, so that its upstream end was above the point of the jet through which H,O, vapour was added, as shown in fig. 1. The inner surface of the liner was coated with a thin layer of boric acid as before,6 to provide the surface upon which the initiation reaction (1) occurred. The components of the reaction mixtures were added as follows. Nitrogen (B.O.C., White Spot) was taken from a cylinder and passed through soda-lime-packed columns to remove traces of CO,. In the experiments at total pressures of 40.0 and 93.3 kPa, N, was used as part of the main carrier gas and in all experiments it was used to carry H,O, and NO, into the main flow.Carbon monoxide was taken from a B.O.C. Technical Grade cylinder and was passed through soda-lime-packed columns to reduce traces of CO, to below the limits of detectability. At 13.3 kPa, CO itself was used as the carrier gas while at 40.0 and 93.3 kPa the carrier gases were 30% CO in N, and 14% CO in N,, respectively: this resulted in a constant [CO] at all pressures. The NO, flow was added to the carrier gas before it entered the main flowtube. A small flowrate of the N, was bubbled through liquified nitrogen dioxide in a series of three bubblers maintained at a temperature of 263 & 1 K using a salt-water-ice freezing mixture.The nitrogen dioxide was taken from a Matheson lecture bottle and was purified as before.8 The rate of addition of NO, to the flow system was monitored by the on-line spectrophotometric system described previously .8 In a similar manner a small flowrate of N, was bubbled through a set of three bubblers containing ca. 100% H,O, liquid, maintained at a temperature of 293 f 1 K. The hydrogen peroxide was obtained as an 85% aqueous solution (Laporte Industries, High Test Peroxide) and was concentrated as described previously.2v The N, + H,O, mixture was added directlyG. J. AUDLEY, D. L. BAULCH A N D I. M. CAMPBELL 1239 to the flowtube through the jet shown in fig. 1. The rate of addition of H,O, was measured by diversion of this flow through a set of bubble chambers containing acidified ceric sulphate solution, with spectrophotometric determination of the resultant rate of removal of CeIV as before.s The rate of flow of N, through the hydrogen peroxide was maintained at 2.3 pmol s-' in all experiments.The two substances which were used in this work as internal kinetic calibrants were acetaldehyde (B.D.H., 99% stated purity) and diethylether (Koch-Light, 99.9% stated purity). / S FIG. 1.-Diagram of the working section of the flowtube. A, circulating water maintained at 298 K; B, inner tube internally coated with boric acid; C, to spiral gauge. These liquids were degassed by pumping, before the vapour from that being used was introduced into an evacuated bulb of volume ca.5 dm3 to a pressure of ca. 5 kPa: N, was then added to bring the total pressure to ca. 110 kPa and mixing was assumed to be achieved on a time scale of at least overnight. Flows from the bulb were added to the main flow when required prior to its entry to the main flowtube. All gas flowrates were measured by capillary flowmeters, calibrated as described before.e At the outlet of the flowtube the reacted gas mixture passed through a sampling bulb of volume ca. 2.2 dm3. When this bulb was considered to contain a sample of gas representative of the conditions in the flowtube, it was isolated from the flow using a by-pass system and the contents were pumped slowly through a spiral trap packed with glass helices maintained at a temperature of 77 K using liquid nitrogen.The resultant condensate was analysed for its CO, content using the gas chromatographic procedures described previously.e Pressures in the flowtube were measured using a calibrated spiral gauge. RESULTS 1 K. All experiments were conducted at 298 The amount of CO, collected in the sample bulb was converted into the concentration of CO,, [CO,], in the reacted mixture on the basis of prior calibrations with known mixtures of CO, in the carrier gas. The activity of the boric-acid-coated surface of the liner tube was determined from measurements of the variation of [CO,] as a function of [S]/[CO] ( S = acetaldehyde or diethyl ether), when rates of addition of NO, and H,O, and total pressure were maintained constant. During the course of this work we encountered three different surface activities, each being stable for a period of months before an unpredictable and virtually discontinuous rise to a higher activity took place.These surface activities are denoted as 1, 2 and 3 in increasing order of efficiency. Plots of the reciprocal of1 240 8.0- 7.0- 6.0- - H,O,+NO,+CO+N, CHAIN REACTION SYSTEM 6 Y' t-- I I 1 I I I I 1 0 2.0 4.0 6.0 8.0 10.0 12.0 103 [SI /[COI FIG. 2.-Plots of the reciprocal yield of CO,, [CO,]-', against the concentration ratio [S]/[CO]. 0, S = diethyl ether, P = 13.3 kPa, surface activity 3; 0, S = acetaldehyde, P = 40.0 kPa, surface activity 1 ; V, S = acetaldehyde, P = 40.0 kPa, surface activity 2; 0, S = diethyl ether, P = 40.0 kPa, surface activity 3; A, S = diethyl ether, P = 93.3 kPa, surface activity 3.A J 0 0.5 1.0 1.5 2.0 2 . 5 1O2"021 ,"COl FIG. 3.-Plots of the reciprocal yield of CO,, [COJ-', against the concentration ratio [NO,]/[CO]. 0, P = 13.3 kPa, surface activity 3; 0, P = 40.0 kPa, surface activity 1; V, P = 40.0 kPa, surface activity 2; 0, P = 40.0 kPa, surface activity 3; A, P = 93.3 kPa, surface activity 3.G. J. AUDLEY, D. L. BAULCH AND I. M. CAMPBELL 1241 [CO,] against fS]/[CO] in the case of each surface activity were found to be good straight lines, as shown in fig. 2 for conditions of total pressure of 40.0 kPa with ca. 30% CO in N, as the carrier gas. For each surface activity we conducted experiments in the absence of added S, with constant input of H,O,, in a carrier gas of ca.30% CO in N, at a constant total pressure of 40.0 kPa, but varying the input of NO, in its excess r e g i ~ n . ~ Plots of [CO,]-' against [NO,]/[CO] proved to be good straight lines as is shown in fig. 3. TABLE G GRADIENTS (G) OF PLOTS OF FIG. 2 AND 3 Gs GNOr P/kPa surface activity S /lo7 dm3 mol-l /lo7 dm3 mol-l 13.3 3 C2H50C2H5 1.16f0.13 0.46 f 0.04 40.0 1 CH3CH0 18.7 f 0.8 4.02 f 0.38 40.0 2 CH,CHO 6.63 f 0.60 2.02 f 0.08 40.0 3 C2H50C2H5 2.52 f 0.13 1.76 & 0.05 93.3 3 C2H2OC2H5 5.01 f0.57 5.38 f 0.21 Also shown in fig. 2 and 3 are more limited data for experiments conducted at total pressures of 13.3 kPa (CO is the main carrier gas) and 93.3 kPa (14% CO in N, is main carrier gas). These data were only obtained for surface activity 3 since surface activities 1 and 2 had disappeared before we undertook the study of the effect of variation of the total pressure.A constant molar flowrate of N, through the hydrogen peroxide bubbler system was used for the addition of an initial concentration of H,O,, denoted as [H,O,],, to the flowtube (which would correspond to the homogeneous concentration of H,O, in the flowtube if no reaction occurred and adsorption on the walls was insignificant). For the total pressures of 13.3,40.0 and 93.3 kPa, [H20210 corresponded to fractions of ca. 1/13, ca. 1/38 and ca. 1/90, respectively, of the saturated vapour pressure of pure hydrogen peroxide at 298 K. The total flowrates were such that typically the time required for the gases to pass down the length of the tube was ca. 80 s, while the half-life for the initial reaction (1) under the conditions used is believed to be at least an order of magnitude ~horter.~ That the reaction time scale was much shorter than the residence time was confirmed by the fact that halving the total flowrate in a number of experiments resulted in no significant variation in [CO,].Least-mean-squares analysis of the plots shown in fig. 2 and 3 yielded gradients G shown in table 1. Those derived from fig. 2 are denoted as G,, when S = aFetaldehyde (CH,CHO) or diethyl ether (C,H,OC,H,) was present, while those derived from fig. 3 are denoted as GN02, when no S was present. The error limits quoted represent one standard deviation. DISCUSSION Analysis of the chain reaction mechanism represented by steps (1)-(4) H,O, +NO, + OH + HNO, OH+CO + CO,+H H+NO,+OH+NO OH +NO, + HNO,1242 H,O,+NO,+CO+N, CHAIN REACTION SYSTEM has been shown2v3 to lead to the equation where [OH], denotes the effective yield of hydroxyl radicals from the initiation step (l), integrated throughout the course of the reaction, expressed as a concentration and [NO,] represents the actual concentration of NO, in the reaction zone: since NO, is added in substantial excess of hydrogen peroxide, [NO,] is always assumed to be equal to the concentration of NO, as added.Provided that the efficiency of the boric- acid-coated surface for conversion of H,O, into OH by reaction (1) remains constant, in a set of experiments where [NO,] is varied at constant total pressure and [H,O,],, eqn (i) predicts that plots of [CO,]-l against [NO,]/[CO] will be linear with zero intercept.Thus the gradients in fig. 3 are expressed by the equation (ii) Upon addition of a substrate S which reacts with OH according to OH + S -+ non-propagating products (9 the analysis2* leads to the equation (iii) In a set of experiments where only [S] is varied, plots of [CO,]-l against [S]/[CO] are predicted to be linear with a positive ordinate intercept and gradient Gs given by When [OH], is constant (i.e. same [H,O,], and same surface efficiency) and when for the substrate used there are available absolute values of k, from other work in the literature, the value of the product k,[OH], may be deduced. Also eqn (ii) and (iv) combine to eliminate k,[OH], so that with corresponding values of GNOt and G, for the same conditions and an absolute value of ks, the apparent value of k4 (ki) can be deduced.For S = CH3CH0, an absolute value of k, = (9.6+ 1.0) x lo9 dm3 mol-1 s-l [ref. (9)] is available and there is also a value of ks = (9.2 f 1 .O) x lo9 dm3 mol-l s-l derived from relative rate measurements.10 Accordingly we have used a mean value of k, = 9.4 x lo9 dm3 mol-l s-l for S = CH3CH0 in this work. For S = C,H,OC,H, there is only one absolute value of k, = (5.4f 1.1) x lo9 dm3 mol-1 s-l available in the literature." However in previous works we have measured the ratio of the G, value for S = C,H,OC,H, to that for S = CH3CH0 as (3.33f0.12)/(6.63 k0.59) = 0.502+0.068. On the basis of the mean value of k, = 9.4 x lo9 dm3 mol-l s-' above this ratio leads to k, = (4.7 k0.6) x lo9 dm3 mol-l s-l for S = C,H,OC,H,, which is consistent with the absolute value.We have therefore used k, = 5.4 x lo9 dm3 mol-1 s-l for S = C,H,OC,H, as the mean value. The values of k2[OH], and k; obtained from the data of table 1 are shown in table 2.G. J. AUDLEY, D. L. BAULCH AND I. M. CAMPBELL 1243 The essential parameter of present interest is [OH],, with regard to establishing the surface efficiency for conversion of [H,O,], in reaction (1). The rate constant k,, for N, carrier gas, is considered generally to increase almost linearly with total pressure (P) so that it is enhanced by a factor of ca. 2 for P x 0.1 MPa compared with the low-pressure limiting value.12v13 As we have postulated b e f ~ r e , ~ the value of k, at our lowest value of P of 13.3 kPa is likely to be indistinguishable, within the TABLE 2.-APPARENT VALUES OF k, AND k2 [OH], surface k; P/kPa [H20,],/10-6 mol dm-3 activity k2[OH],/102 s-l /lo9 dm3 mol-l s-l 13.3 6.40 f 0.45 3 4.7 + 0.5 2.14 & 0.29 40.0 2.36 f 0.22 1 0.50 f 0.02 2.02 f 0.21 40.0 2.36 f 0.22 2 1.4kO.l 2.86 f 0.28 40.0 2.36 f 0.22 3 2.1 f O .1 3.77 0.23 93.3 0.83 +_ 0.06 3 1 . 1 +O.l 5.80 f 0.70 experimental error limits, from the low-pressure limiting value of k, = (9.0+0.9) x lo7 dm3 mol-l s-l. On this basis, for P = 13.3 kPa and surface activity 3, the value of k,[OH], shown in table 2 yields [OH], = (5.2k0.6) x mol dm-3 when [H,O,], = (6.40 f 0.45) x mol dm-3, corresponding to a surface conversion efficiency for reaction (1) of (8 1 lo)%. If k, were assumed to be independent of P in our experiments, then the values of k,[OH], shown in table 2 would indicate values of the surface conversion efficiency rising to well over 100% at the higher pressures.This appears most unlikely. On the other hand, if we assume that the surface conversion efficiency is independent of P and calculate corresponding values of k,, at P = 40.0 and 93.3 kPa we obtain sensible results in comparison with those from other studies in which the pressure-dependence of k, has been investigated in N, diluent. For example, using a surface conversion efficiency of (81 & lo)% in conjunction with the k,[OH], shown in table 2 for P = 93.3 kPa, we obtain k, = (1.6 f 0.3) x lo8 dm3 mol-l s-l, which is enhanced by the expected factor of ca. 2 compared with the low-pressure limiting value of k,.The fact that the carrier gas is composed of CO as well as N, is not expected to be of significance in this: evidence is presented below in connection with k, that CO and N, have approximately the same third-body efficiencies in three-body combination reactions. We therefore conclude that the most satisfactory interpretation of our results is based on a pressure-independent surface efficiency of reaction (1) and a pressure-dependent k,. Biermann et all3 have shown that k, in N, carriers only shows a pressure-dependent component when traces of 0, (ca. 0.03 % of [N,] for P x 0.1 MPa) or other (undefined) impurities are present to react with the M-stabilized HOCO complex involved. However, in this work we have taken no specific precautions for removal of 0, and it might not be unreasonable to suspect that NO, may be effective in the same role.Accordingly our interpretation that k, shows a pressure dependence in H,O, + NO, + CO systems is not surprising. Turning to reaction (4), it is immediately clear from the three sets of data for P = 40.0 kPa that k; is not independent of the surface conversion efficiency of reaction (1): k; increases as k,[OH], (i.e. [OH],) increases. These values are for N, carrier containing 30% CO but when an N, carrier containing only ca. 2.5% CO was used no significant difference ink; was found, confirming that N, and CO have approximately the same third-body efficiency in reaction (4). Anastasi and Smith14 have obtained1244 H,O,+NO,+CO+N, CHAIN REACTION SYSTEM absolute values of k, as a function of P for M = N, at ambient temperature, which have been closely confirmed by Wine et aZ.,15 from which we interpolate k, = (4.4 & 0.8) x lo9 dm3 mol-l s-l, using the 18% error limit quoted by Anastasi and Smith.’, In table 2 the values of ki for P = 40.0 kPa are significantly lower than the above value of k,, but tend towards it as the surface conversion efficiency increases.Assuming the mean [H,O,], to [OH], conversion efficiency of 8 1 % for surface activity 3, then surface 1 corresponds to a conversion efficiency of ca. 20% and surface activity 2 to ca. 53 %. A plot of ki against surface conversion efficiency for the values obtained at P = 40.0 kPa is approximately linear as is shown in fig. 4. 8.0 7.0 6.0 - 5.0.I 4 E “E 4.0. 2 a 01 + 1.0 0 ; I I I 1 I 0 20 40 60 80 100 surface efficiency (%) FIG. 4.-Plot of ki against surface efficiency. 0, P = 13.3 kPa; 0, P = 40.0 kPa; A, P = 93.3 kPa. The extrapolation of the linear trend to 100% conversion efficiency leads to a value of ki which is in agreement, within relatively wide error limits, with the absolute value of k, for P = 40.0 kPa quoted above. In fact this is the situation which we found in our previous work3 in a static reactor in which the surface conversion efficiency was 100%: apparent values of k , extracted were within experimental error of the absolute values of k , derived from Anastasi and Smith’s results14 for P < 13.3 kPa. Also indicated in fig. 4 is that the two values of ki for P = 13.3 kPa and P = 93.3 kPa for surface activity 3 tend similarly to be lower than the corresponding values of k, derived from Anastasi and Smith’s work.l4 The interpretation of low values of ki compared with k, seems most likely to originate from a physical rather than a chemical effect in the flowtube, stemming from the mixed heterogeneous/homogeneous nature of the H,O, + NO, + CO reaction system.The heterogeneous initiation step (1) generates OH radicals at the boric- acid-coated surface and these diffuse out into the gas phase. In our present experiments a typical chain length is 4 and, since k , is approximately three orders of magnitudeG. J. AUDLEY, D . L. BAULCH A N D I. M. CAMPBELL 1245 less than the collision frequency between CO and OH, the average OH will make ca.lo3 collisions with CO molecules before it reacts and is converted to an H atom. Since the rates of the propagation steps (2) and (3) are equal, the average H atom survives for the same time as the OH radical before it reacts with NO, to be reconverted to OH. On average this cycle can be considered to occur 4 times before the chain propagation centre is extinguished by the termination reaction (4). As we have argued b e f ~ r e , ~ the chain propagation cannot extend to a distance far from the liner wall in comparison with the tube radius, despite the promotion of OH and H diffusion by their respective concentration gradients from the wall. Thus it is necessary that the entire reaction takes place in a thin layer of the gas phase adjacent to the liner tube wall, which contains instantaneously only a very small fraction of the NO, present in the flowtube.If initiation becomes sufficiently rapid, NO, consumption in reactions (1) and (4) and its conversion to NO in reaction (3) could be imagined to overwhelm the diffusional ability of NO, to maintain the bulk gas-phase [NO,] across this reaction zone, Accordingly a concentration gradient of [NO,] would develop in this thin layer towards the wall. Under such circumstances the average [NO,] in the reaction zone would be less than that in the bulk gas phase: this would reduce the termination rate and hence the apparent value of k, deduced under the assumption that [NO,] in the bulk gas phase was applicable to the reaction zone. On the other hand, the actual chain length in the reaction zone, dependent on the ratio k,[CO]/k,[NO,], will increase, offsetting the reduction in initiation rate to match the reduced termination rate.NO, formed in reaction (3), will be a component of the reaction zone : however, the rate constant of the additional termination reaction OH +NO( +M) + HNO,( +M) ( 5 ) is approximately a factor of 3 lower than k, [ref. (16)] under our conditions. Effective conversion of NO, to NO in the boundary layer, exacerbated by the increased chain length, therefore leads to ki < k, in our analysis. CO is always present in such a large excess that a significant gradient of [CO] across the reaction zone is unlikely. Added S molecules are only consumed at a fraction of the termination rate, so that significant S depletion in the reaction zone is also unlikely.On the basis of our model, depletion in the boundary layer will be most developed above surface 1, when the value of k; shown in fig. 4 suggests an effective [NO,] depletion of ca. 50%. Depletion of S in the boundary layer under these extreme circumstances is estimated to be limited to ca. 10% for the range of [S] used in this work and is therefore insignificant for the general conclusions drawn in connection with surface 1. For the higher efficiencies represented by surfaces 2 and 3, depletion of [NO,] and hence [S] in the boundary layer will be less. On the basis of the above model, our experimental observation that ki increases with improving surface efficiency points to a more severe depletion of [NO,] in the reaction zone at lower conversion efficiencies of [H,O,], into [OH],.The additional factor in the system which is required to resolve this apparent anomaly is the activity of the surface for H,02 adsorption, which is the pre-requisite for reaction (1) according to the earlier work of Gray et a1.l’ If the surface is very active in this respect, the rate of adsorption and the surface coverage will be large: the initially added H,O, will then diffuse to the liner wall to produce a high surface concentration of H,O, on a ring of the surface with a small lengthwise spread. The adjacent reaction zone will have a small volume under these circumstances with a high production rate of OH from reaction (1) per unit volume. This situation has an evident potential for depletion of [NO,] in the highly localized reaction zone.Moreover, such a high effective concen- tration of H,O, may promote interaction of the nascent OH with adsorbed H202 at1246 H,O,+NO,+CO+N, CHAIN REACTION SYSTEM the expense of interaction with CO in the gas phase, thus accounting for the inefficiency of [H202], to [OH], conversion. An alternative explanation is that reaction (1) proceeds very rapidly on this type of surface, which produces a similarly localized reaction zone but cannot account so readily for the inefficiency of conversion. The surface with higher [H,O,], to [OH], conversion efficiencies are then seen as less active for adsorption of H,O,. The ring of surface upon which the initially added H,O, is adsorbed extends much further parallel to the axis of the tube.Reversal of the arguments in the preceding paragraph suggests that due to the larger volume of the reaction zone and the lower production rate of OH per unit volume from reaction (l), the potentials for depletion of [NO,] and for interaction of OH with adsorbed H,O, are considerably reduced. Consequently kl and the [H,O,], to [OH], conversion efficiency are both increased. Although this model of the reaction zone is qualitative and could only be substantiated by a much more sophisticated technique, it provides a reasonable interpretation for all the results presently available for the H,O, + NO, + CO reaction system. The sequence of events in our work is then that the initial surface activity 1 was induced by a surface with a high activity for adsorption of H,O,.Discontinuously, for unknown reasons, the surface then deactivated for H,O, adsorption to produce surface activity 2. Finally a further transition, with further deactivation, produced surface activity 3. Deactivation for adsorption over a long period of operation, as demanded by these arguments, seems more in keeping with the general experience with catalytic surfaces than the activation which might at first sight be ascribed to the observed increase of [H,O,], to [OH], conversion efficiency with time. Finally we consider the likelihood that reaction (1) is more complex than is indicated by the reaction equation. Hydrogen peroxide in aqueous solutions decomposes in the presence of mixed insulator oxide materials possessing acidic centres. A mechanism which has been proposedlB for the decomposition is A++H,O, + A+'HO;+H+ (4 A+'HO;+H,O, + A+'OH-+H,O+O, (b) A+'OH-+H+ +A++H,O (4 in which AS represents an acidic centre.Whilst our H,O, + NO, interaction takes place at the boundary between solid and gas phases, a modification of step (b) in the above mechanism appears to offer some rationalization. Boric acid will evidently have acidic centres which can be represented in general function as A+. In their original work on the H20,+N0, systems, Gray et aZ." suggested that an adsorbed H,O, molecule interacted with a gas-phase NO, molecule. Adapting step (b) above we suggest a step A+'HO;+NO, +A+'NO;+OH in which the A+'NO; species does not appear unrealistic and like step (b) atom transfer is involved. Finally the process is completed by the step represented as A+'NO; + H+ + A+ + HNO,.Further work is required to justify such a mechanism, which must be regarded as speculative for the present. One of us (G.J.A.) thanks the S.R.C. for the award of a CASE studentship. We thank Laporte Industries Ltd for gifts of hydrogen peroxide solution and for their interest in the work.G. J. AUDLEY, D . L. BAULCH A N D I. M. CAMPBELL 1247 D. L. Baulch and I. M. Campbell, Gas Kinetics and Energy Transfer (Specialist Periodical Reports, Royal Society of Chemistry, London, 1981), vol. 4, pp. 137-188. I. M. Campbell, B. J. Handy and R. M. Kirby, J. Chem. SOC., Faraday Trans. I , 1975,71, 867. I. M. Campbell and P. E. Parkinson, J. Chem. SOC., Faraday Trans. I , 1979, 75, 2048. I. M. Campbell, D. F. McLaughlin and B. J. Handy, Chem. Phys. Lett., 1976,38, 362. I. M. Campbell and P. E. Parkinson, Chem. Phys. Lett., 1978, 53, 385. G. J. Audley, D. L. Baulch and I. M. Campbell, J. Chem. SOC., Faraday Trans. I , 1981, 77,2541. ' I. M. Campbell and K. Goodman, Chem. Phys. Lett., 1975,36, 382. * G. J. Audley, D. L. Baulch, I. M. Campbell, D. J. Waters and G. Watling, J. Chem. SOC., Faraday Trans. I , 1982, 78, 611. R. Atkinson and J. N. Pitts, J. Chem. Phys., 1978, 68, 3581. lo H. Niki, P. D. Maker, C. M. Savage and L. P. Breitenbach, J. Phys. Chem., 1978, 82, 132. A. C. Lloyd, K. R. Darnall, A. M. Winer and J. N. Pitts, Chem. Phys. Lett., 1976, 42, 205. l2 R. A. Cox, R. G. Dement and P. M. Holt, J. Chem. SOC., Faraday Trans. I , 1976, 72, 2031. l3 H. W. Biermann, C. Zetsch and F. Stuhl, Ber. Bunsenges. Phys. Chem., 1978, 82, 633. l4 C. Anastasi and I. W. M.-Smith, J. Chem. SOC., Faraday Trans. 2, 1976, 72, 1459. l5 P. H. Wine, N. M. Kreutter and A. R. Ravishankara, J. Phys. Chem., 1979, 83, 3191. l6 C. Anastasi and I. W. M. Smith, J. Chem. SOC., Faraday Trans. 2, 1978, 74, 1056. D. Gray, E. Lissi and J. Heicklen, J. Phys. Chem., 1972, 76, 1919. S. P. Walvekar and A. B. Halgeri, 2. Anorg. Allg. Chem., 1973, 400, 83. (PAPER 1/942)

ISSN:0300-9599    

DOI:10.1039/F19827801237

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

Ultrasonic Relaxation in Aqueous Solutions of Butanediols BY SADAKATSU NISHIKAWA* AND MITSUO MASHIMA Department of Chemistry, Faculty of Science and Engineering, Saga University, Saga 840, Japan Received 12th June, 1981 Ultrasonic absorption and velocity have been measured to investigate the structural and dynamic properties of aqueous solutions of butane-1,Cdiol and butane-1,2-diol at 20 OC. In the former solution no excess absorption was found, and in the latter solution a single relaxational process was observed in the frequency range 15-220 MHz. The excess absorption mechanism has been analysed as a solute-solvent interaction, AB A+ B, and the rate constants for butane- 1 ,2-diol solution have been determined to be kf = 1.5 x 108 s-' and kb = 1.6 x 108 dm3 mo1-l s-' for the forward and backward steps, respectively.The influence of the position of hydroxy groups in the molecules on the water structure is discussed by a comparison with the results for solutions of the two dihydric alcohol. In preceeding papers'v we have reported the structure and kinetics of aqueous solutions of alcohols and ethers by means of ultrasonic methods. The characteristic properties of these solutions are the peak sound velocity concentration and the peak sound absorption concentration. Mechanistic models3 of these properties have been proposed for various non-electrolyte solutions, and more recently fluctuation models4~ have been suggested. Our interpretation of the absorption mechanism is based on an analysis of the concentration dependences of the relaxation frequencies; this is related to reaction rates and the maximum excess absorption per wavelength, which is associated with the volume and enthalpy changes of the reaction.We also base our interpretation on the nature of the solute molecules. The ultrasonic studies have been extended to aqueous solutions of dihydric alcohols in order to investigate whether the proposed model for aqueous solutions of alcohols and ethers is also applicable. For this purpose we have chosen butane- 1,2-diol and butane- 1,4-diol as the solutes. EXPERIMENTAL Butane- 1,2-diol and butane- 1 ,Cdiol were obtained from the Tokyo Kasei Kogyo Co., and were distilled under reduced pressure. The chemicals were subjected to gas-chromatographic analysis and showed purities of < 99.8%.Doubly distilled water was used as solvent. The desired concentrations were obtained by weighing. The ultrasonic pulse techniques was used to measure the absorption coefficient, a, in the frequency range 15-95 MHz using a 5 MHz fundamental crystal and in the range 60-220 MHz using a 20 MHz crystal. The uncertainty in the absorption coefficient was within f: 2.5 %. The ultrasonic velocity was measured by an interferometer operating at 2.5 MHz; uncertainties in the measurement were within f: 50 cm 0. The solution density was measured using a standard pyknometer having a volume of ca. 2.3 cm3. All the measurement cells were immersed in a water bath which was maintained at a constant temperature to better than fO.O1 OC. In order to calculate the ultrasonic, kinetic and thermodynamic parameters a Hitachi microcomputer was used.41 1249 FAR 11250 ULTRASONIC RELAXATION I N AQUEOUS BUTANEDIOLS RESULTS No frequency dependence of a / f 2 , wherefis the frequency, has been observed in aqueous solutions of butane-1,4-diol up to a concentration of 6.8 mol dm-3. On the other hand, excess absorption was found in solutions of butane- 1,2-diol. The observed spectra are all characteristic of a single relaxation, which is expressed as follows: or where A is the amplitude of the excess absorption, B is that of the background absorption, fr is the relaxation frequency, p is the excess absorption per wavelength, and c is the velocity of sound. Fig. 1 shows representative absorption spectra in aqueous solutions of butane- 1,2-diol and butane- 1,4-diol, which are expressed in terms of eqn (1).The ultrasonic parameters were determined so as to obtain the best fit to 1 I I 0 2o 10 50 100 500 fIMHz FIG. 1 .-Representative ultrasonic absorption spectra for aqueous solutions of butane- 1,2-diol and butane-lP-diol at 20 O C . Arrows show the relaxation frequency and solid curves are calculated using eqn (1). Butane-1,2-diol: (>, 2.25; 0, 3.04; 0, 5.07mol dm-3. Butane-1,Cdiol: 0, 2.48; 0, 4.13; 0, 6.82 mol dm-3. TABLE 1 .-ULTRASONIC PARAMETERS FOR AQUEOUS SOLUTION OF BUTANE- 1,2-DIOL 4 C e A B /mol dm-3 /10-17 s2 cm-l /10-17 s2 cm-l fr/MHz P / g c/m s-I 2.25 35.5 19.7 250 1.011 1627 3.04 46.6 48.4 173 1.015 1654 4.12 109 113 131 1.020 1667 5.07 169 131 114 1.020 1659 6.18 128 - 204 113 1.022 1638 6.90 105 220 140 1.021 1628 7.95 95.1 260 130 1.018 1594S.NISHIKAWA AND M. MASHIMA 1251 I I I 1 3 5 7 CJmol dm-3 FIG. 2.4oncentration dependences of excess absorption, A , and the background absorption, B. a, B for butane-1,2-diol; 0, A for butane-1,2-diol; 0, Bfor butane-1,Cdiol; 0, B for butane-1,Cdiol measured by Blandamer et al.' the experimental data, using a least-mean-squares method. Those for butane- 1,2-diol are given in table 1; the values of A and B are shown in fig. 2, along with those of B for butane- 1,4-diol as a function of concentration. In both solutions, the background absorption increases with concentration and has a higher value than would be expected from classical absorption, i.e. absorption due to viscous and thermal effects; this means that other excess absorptions are predicted in the higher frequency range.The amplitude of the excess absorption, A , for butane-1,2-diol shows a maximum at ca. 5.2 mol dm-3. The sound velocity passes through a maximum at 4.0 mol dm-3. Similar dependences of concentration on ultrasonic properties have been observed in aqueous solutions of monohydric alcohols and The proposed modell for the excess absorption mechanism in the megahertz frequency range is the interaction between solute and solvent, in which only non-hydrogen-bonded water molecules may participate. This is shown in reaction (2) where AB is the complex, A is the solute, B is the non-hydrogen-bonded water molecule and Ci are the equilibrium concentrations of each component. An important factor which may arise in the analysis of the relaxational process is a coupling to other reactions which occur in solution, e.g.internal rotation of the diols. When the relaxation frequencies are relatively close to each other, the slower process should be 41-21252 ULTRASONIC RELAXATION I N AQUEOUS BUTANEDIOLS influenced by the faster one. However, in aqueous solutions of butane- 1,2-diol, only a single relaxation process has been observed, and the influence caused by other reactions in our frequency range may be so small that, to a good approximation, only one equilibrium perturbation is the cause of the excess absorption. The relation between the relaxation frequency and the analytical concentrations of the solute and solvent is derived as follows for reaction (2): where Ce, Cw, B and K are the analytical concentrations of the solute and solvent, the mole fraction of non-hydrogen-bonded water and the equilibrium constant defined by K = k,/k,, respectively.The relaxation frequency passes through a minimum at 5.2 mol dm-3, and this condition may be used to obtain the relationship between j? and K. Then kb, B and K values are determined so as to obtain the best fit to eqn (3). The calculated values using these determined parameters are shown in fig. 3 by a solid line. In table 2 the rate and thermodynamic constants are listed along with those for monohydric alcohol solutions for comparison. 3001 ; I I I 1 I 3 5 7 9 CJmol dm-3 FIG. 3.-Relaxation frequency as a function of butane-1,2-diol concentration. The solid curve was calculated for the mechanism AB e A+ B, and the dashed one was determined for B = 0.175 and K = 0.620 dms m o P for the mechanism AB, e A + 2 B .TABLE 2.-RATE AND THERMODYNAMIC CONSTANTS FOR ALCOHOLS AND DIOL k b kf/s-' /dm3 mol-l s-l B ref. - 0.26 8 pure water - isopropyl alcohol 1.4 x lo8 8.8 x lo7 0.17 lc butane- 1,2-diol 1.5 x lo8 1 . 6 ~ lo8 0.19 this work propyl alcohol 1.6 x lo8 6.2 x 107 0.15 la t-butyl alcohol 1.2 x lo8 6.3 x 107 0.12 lbS. NISHIKAWA A N D M. MASHIMA 1253 Another parameter determined from ultrasonic absorption measurements is the maximum excess absorption per wavelength, pmax. For reaction (2) we have Af,c 7tpc2r - (AV- up AH/pCJ RT Pmax = - - 2 (4) r = ( i / v ) ( i / c l + 1/c2+ 1/c3- i/cT)-l ( 5 ) where p is the density, R is the gas constant, T is the absolute temperature, A V is the volume change occurring during the reaction, AH is the enthalpy change, up is the thermal expansion coefficient, C, is the specific heat at constant pressure, and V is the molar volume.To a good approximation pmax is almost proportional to p c T V , because the contribution of the concentration dependences of other terms to pmax is very small. Fig. 4 shows plots of pmax and p c T V . These confirm that the excess absorption mechanism may be due to the perturbation of the equilibrium shown in reaction (2). 2 I I I I 3 5 7 C,/mol dm-3 FIG. 4.4oncentration dependences of p,,, and p c T V for aqueous solutions of butane-1 ,Zdiol. DISCUSSION Andreare et aL3 have considered and reviewed various mechanisms in order to interpret the characteristic properties of non-electrolytes in aqueous solution.Recently approaches to the absorption mechanism have been taken into account using fluctuation However, depending upon the choice of solute, the magnitudes of the peak sound absorption are different and one or two relaxational absorptions are observed. Therefore we prefer to assume the normal perturbation of equilibrium as the cause of the excess absorption mechanism. Blandamer et a1.' have reported ultrasonic absorption results in butane- 1,4-diol aqueous solution in which no excess absorption was observed. They also found that the existance of the excess absorption in the megahertz frequency range depends on the structure of the alcohol, which seems to be consistent with our results.1254 ULTRASONIC RELAXATION I N AQUEOUS BUTANEDIOLS Although the interaction between butane-1,2-diol and water is attributed to the formation of a 1 : 1 complex, it might be expected that more than one water molecule is bonded to butane-1,2-diol because of the Occurrence of two hydroxy groups in the molecule.We also tested to see if the AB, e A + 2B mechanism was associated with the excess absorption. The relaxation frequency for this mechanism is expressed as 2nfr = kk(4[A] [B] + [BI2) + k;. Assuming appropriate values of /? and K (the detailed procedure for this calculation is described in a previous papere), each equilibrium concentration was calculated and the ratio k;/kk was determined so as to obtain the same value of Kas had been assumed and so as to reach a maximum of pcT’V when r’ = (1 / C , + 4/C, + 1 /C3 - 4/CT)/ V.When /? values in the range 0.16-0.19 are chosen, a K value which satisfies the above conditions is obtainable. However, the concentration dependence of the relaxation frequency is not so well interpreted (one set of results is shown in fig. 3 by a dashed line), and the probable errors in k; and k6 were larger than the probable values. Another reason for our conclusions is that the observed excess absorption is characteristic of a single relaxation process and therefore only one equilibrium perturbation between different molecular species is the cause of the observed excess absorption. One possible explanation of the ,single relaxation might be that intra- molecular hydrogen bonding could exist in the butane-l,2-diol molecule because of the very close positions of the two hydroxy groups.It is thus strange that no excess 100- 50 - I F - I- - I 2 3 4 5 6 carbon no./no. of OH groups FIG. 5.-Plots of the relaxation frequency range and carbon number per hydroxy group. 0 Butane-1,2-diol, 0 ethyl alcohol, 0 isopropyl alcohol, @ propyl alcohol, 0 t-butyl alcohol, 8 propyl cellulose, 0 butyl cellulose. absorption was found in butane- 1,4-diol solutions. By examining the ultrasonic absorption data for aqueous solution of various alcohols and ethers reported so far, it is found that the excess absorption depends strongly on the size of the alkyl group. In fig. 5 , the ranges of relaxation frequency observed in the solutions are plotted against the carbon number per hydroxy group.As the size of the hydrophobic group increases, the excess absorption appears in a lower frequency range and in a lower concentration range. Butane-l$diol has two hydroxy groups at each end and it may be considered that the number of alkyl groups per hydroxy group is two; however, the effect of the size of the alkyl group on the hydroxy group is less than that forS. NISHIKAWA AND M. MASHIMA 1255 butane-l,2-diol or ethyl alcohol. The excess absorption in ethyl alcohol solution is observed over a high frequency range,l0 and no excess absorption in methyl alcohol solution has been reported in the megahertz frequency range at room temperature.ll Therefore the hydrophobicity of butane- 1,4-diol seems to be small, and less than that of butane-l,2-diol and ethyl alcohol. At this stage we may conclude that the interaction between water and solutes containing hydroxy groups is strongly dependent on the strength of the hydrophobicity of the solute; the /3 parameter is very useful in estimating the effect of the solute on the water structure.We thank Miss F. Ikegami for performing some of the ultrasonic measurements. (a) S. Nishikawa, M. Mashima and T. Yasunaga, Bull. Chem. SOC. Jpn, 1975,424,661 ; (b) S. Nishikawa, M. Mashima, M. Maekawa and T. Yasunaga, Bull. Chem. SOC. Jpn, 1975,48,2353; (c) S. Nishikawa, M. Mashima and T. Yasunaga, Bull. Chem. SOC. Jpn, 1976, 49, 1413. * S. Nishikawa, M. Tanaka and M. Mashima, J. Phys. Chem., 1981, 85, 686. J. H. Andreae, P. D. Edmonds and J. F. McKellar, Acustica, 1965, 15, 74. (a) G. Atkinson, M. M. Emara, H. Endo and B. L. Atkinson, J. Phys. Chem., 1980, 84, 259; (b) G. Atkinson, S. Rajagopalan and B. L. Atkinson, J. Phys. Chem., 1981, 85, 733. (a) Y. Harada, Y. Suzuki and Y. Ishida, Phys. Rev. A, 1980, 21, 928; (b) Y . Harada, J. Phys. SOC. Jpn, 1980, 48, 705. N. Tatsumoto, J. Chem. Phys., 1967, 47, 4561. 1805. T. A. Litovitz and E. H. Carnevale, J. Appl. Phys., 1955, 26, 816. S. Nishikawa, Y. Yamashita and M. Mashima, Bull. Chem. SOC. Jpn, 1982, 55, in press. ' M. J. Blandamer, N. J. Hidden, M. C. R. Symons and N. C. Treloar, Trans. Faraday SOC., 1969,65, lo K. Takagi and K. Negishi, Jpn J. Phys., 1975, 14, 953. l1 J. Emery and S. Gasse, Acustica, 1979, 43, 205. (PAPER 1/949)

ISSN:0300-9599    

DOI:10.1039/F19827801249

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. 1, 1982, 78, 1257-1267 Standard Potentials of the Silver, Silver Bromide Electrode in Tetrahydrofuran and Tetrahydrofuran + Water Mixtures at Different Temperatures and Related Thermodynamic Quantities BY MAHMOUD M. ELSEMONGY,*~ IBRAHIM M. KENAWY AND ABDELAZIZ S. FOUDA Chemistry Department, Faculty of Science, Mansoura University, Egypt Received 6th June, 198 1 The standard potentials of the Ag,AgBr electrode have been determined in tetrahydrofuran (THF) and in nineteen THF + water solvent mixtures from the e.m.f. measurements of the cell Pt I H,(g, 1 atm)$ I HBr (m), solvent I AgBr (s) I Ag at intervals of 5 OC from 5 to 45 OC. In solvents of high THF content, where the dielectric constant is small, it was necessary to correct for ion-pair formation.The temperature variation of the standard potential has been utilized to evaluate the standard thermodynamic functions for the cell reaction and the standard thermodynamic quantities for the transfer of HBr from water to the respective solvents. The results are discussed in the light of ion-solvent interactions as well as the structural changes of these solvents. Dipolar aprotic solvents are particularly interesting in that they solvate anions much less strongly than do protic solvents such as water.' This is usually reflected in the solvent effects on acid-base strengths, rates of reactions and conductance, for instance. A better understanding of these effects in non-aqueous, as well as partially aqueous, media can be obtained from fundamental electrochemical studies which offer useful information with regard to the nature of ion-solvent interactions.However, relatively few electrochemical studies have been made in aqueous mixtures of dipolar aprotic solvents. Recently we have reported2 the standard potentials of the Ag , AgCl electrode and related thermodynamic quantities in such dipolar aprotic solvent media as tetrahyd- rofuran (THF)+water mixtures containing up to 90% (w/w) THF at 15-55 OC. The Ag , AgBr electrode has distinct advantages over the Ag, AgCl electrode, particularly for the determination of the dissociation constants of nitrogen bases, in which AgBr is less ~oluble.~ Thus, the standard potential of the Ag,AgBr electrode has been determined by Roy et aL3 in three THF+water mixtures containing 10, 30 and 50% (w/w) THF at 5-55 OC.Thermodynamic properties of hydrobromic acid in these solvents have also been rep~rted.~ However, no work seems to have been done on the determination of the standard potentials of the Ag , AgBr electrode in pure THF or its aqueous mixtures other than for these three THF+water solvents. Thus as a part of a comprehensive study on solute-solvent interactions and the t Present address: Chemistry Department, Faculty of Science, Kuwait University, P.O. Box 5969, $ 1 atm = 101 325 Pa. Kuwait . 12571258 STANDARD POTENTIALS OF THE Ag,AgBr ELECTRODE related structure of solvents in both aqueous and non-aqueous media2*4*5 based on e.m.f. measurements of the cell PtIH, (g, 1 atm)lHBr (m), solventlAgBr (s)IAg (1) we have determined the standard potentials of the Ag,AgBr electrode in THF and in nineteen THF + water mixtures at nine different temperatures ranging from 5 to 45 OC.These have been used to evaluate the transfer functions AGP and ASP for HBr in the respective solvents. We can thus obtain information about thermodynamic properties of HBr in these solvents and about the properties and structure of the solvents. EXPERIMENTAL The preparation and analysis of hydrobromic acid solutions in the various solvent mixtures were identical to the methods described previ~usly.~ The molality of the HBr solutions ranged from 0.01 to 0.1 mol kg-l. The acid concentration was accurate to within +0.01%. THF (B.D.H. grade) was purified in the manner described previously.2 The middle fraction of the second distillate was subsequently used in the preparation of the cell solutions.The purity of the sample was verified by gas-liquid chromatography. The purified solvent was used within two days. The aqueous solvents were made by weighing, vacuum corrections being applied to all weighings. The conductivity of water used in the preparation of the solutions was < 0.7 x i2-l cm-l. The THF content of all the solutions reported was accurate to within f 0.01 %. The solutions were all stored in dark bottles shielded from the light under a nitrogen atmosphere. All solutions were swept through with a stream of purified hydrogen gas for 1-2 h before the cells, which were fitted with triple saturators, were filled, to prevent contamination of the solutions with air. All solutions were freshly prepared before taking measurements.The electrodes were prepared essentially as described el~ewhere.~ The experimental set-up and the general procedure used for the e.m.f. measurements were identical with those given previously.4* The measurements were made with three hydrogen electrodes and three Ag, AgBr electrodes for each solution, at intervals of 5 O C from 5 to 45 "C. The cells were thermostatted at each temperature with an accuracy of fO.01 O C . The Ag,AgBr electrodes were found to be stable over the entire temperature range, and the constancy of the cell e.m.f. to f0.05 mV over a period of 1 h was considered as an adequate criterion of equilibrium in the e.m.f. measurements. As a precaution, a given cell was never measured over the entire temperature range.Three series of results were made at each acid concentration. The first was from 5 to 25 O C , the second from 20 to 35 "C and the third from 30 to 45 "C. As new solutions were prepared for the measurement in each, the results serve as an excellent means of checking the reproducibility of the procedure. The e.m.f. values were generally reproducible to f 0.05 mV for different solutions. The cell measurements were in triplicate, and the mean values of these observations recorded. The triplicates generally agreed within f 0.09 mV. RESULTS AND DISCUSSION The measured values of e.m.f. were corrected in the usual way to a pressure of hydrogen of 1 atm. The properties of the THF +water mixtures over the temperature range 5 to 45 *C were derived from previous data.2$ 3$ 8 y The standard potentials of the Ag,AgBr electrode in the water-rich solvents have been determined at each temperature by the usual extrapolation technique, making use of the extended terms of the Debye-Huckel theory, and the procedure is essentially the same as that used in our previous determinatiom29 4 9 In the THF-rich solvents, where the dielectric constant is < 25, ion-pair formation occurs and hydrobromic acid behaves as a weak electrolyte.Thus corrections for ion association in these solvents were taken into account. The standard e.m.f. can no longer be obtained by the simple procedure, but a method that involves preliminary knowledge of the ionization constant of HBr isTABLE 1 .-STANDARD MOLAL POTENTIALS (Eg/V) OF THE Ag , AgBr ELECTRODE IN TETRAHYDROFURAN +WATER SOLVENT MIXTURES AT 5-45 O C ~ ~ ~~ ~~ ~ _ _ _ _ _ _ _ _ ~~ temperature/OC THF (wt %) 5 10 15 20 25 30 35 40 45 0 5 10 15 20 25 30 35 40 45 50 55 60 65 75 80 85 90 95 1 00 70.0.079 62 0.078 07 0.076 79 0.076 08 0.075 75 0.075 64 0.075 08 0.073 16 0.070 22 0.066 33 0.061 60 0.055 85 0.049 44 0.041 91 0.033 47 0.024 26 0.014 49 0.004 37 -0.006 08 -0.016 79 -0.028 02 0.077 73 0.076 21 0.075 08 0.074 19 0.073 65 0.073 18 0.072 32 0.069 84 0.066 50 0.062 11 0.057 07 0.051 07 0.044 33 0.036 35 0.027 78 0.0 18 44 0.008 62 -0.001 59 -0.012 07 -0.023 06 -0.034 13 0.075 67 0.074 39 0.073 19 0.072 24 0.071 52 0.070 71 0.069 29 0.066 36 0.062 58 0.057 82 0.052 44 0.046 03 0.038 85 0.030 78 0.021 82 0.012 23 0.002 28 -0.008 02 -0.018 56 -0.029 32 -0.040 64 0.073 42 0.072 21 0.071 12 0.070 03 0.069 01 0.067 78 0.066 02 0.062 78 0.058 46 0.053 37 0.047 55 0.040 64 0.033 31 0.024 77 0.015 59 0.005 98 -0.004 07 -0.014 50 -0.025 06 -0.035 98 -0.047 32 0.071 05 0.069 98 0.068 89 0.067 70 0.066 43 0.064 77 0.062 42 0.058 76 0.054 13 0.048 56 0.042 29 0.035 14 0.027 39 0.018 67 0.009 28 -0.000 59 -0.010 82 -0.021 31 -0.031 99 -0.042 94 - 0.054 2 1 0.068 52 0.067 49 0.066 50 0.065 16 0.063 61 0.061 54 0.058 59 0.054 62 0.049 45 0.043 58 0.036 83 0.029 32 0.021 24 0.012 26 0.002 62 -0.007 39 -0.017 68 -0.028 33 -0.039 05 - 0.050 16 -0.061 43 0.065 85 0.064 98 0.063 79 0.062 25 0.060 45 0.057 88 0.054 61 0.050 07 0.044 71 0.038 19 0.031 14 0.023 37 0.014 97 0.005 72 -0.004 18 -0.014 33 - 0.024 92 -0.035 57 -0.046 54 -0.057 53 -0.068 89 0.063 02 0.062 11 0.060 97 0.059 33 0.057 13 0.054 22 0.050 35 0.045 38 0.039 62 0.032 80 0.025 25 0.017 02 0.008 31 -0.001 13 -0.011 16 -0.021 69 -0.032 36 -0.043 22 -0.054 12 -0.065 39 - 0.076 8 1 0.060 03 0.059 26 0.057 97 0.056 02 0.053 54 0.050 17 0.045 88 0.040 56 0.034 15 0.026 94 0.019 15 0.010 63 0.001 62 -0.008 21 -0.018 54 -0.029 07 -0.039 93 -0.050 91 -0.062 05 -0.073 24 -0.084 791260 STANDARD POTENTIALS OF THE Ag,AgBr ELECTRODE TABLE VALUES OF THE CONSTANTS a, b AND c OF EQN (1) FOR EVALUATION OF Erne IN THE Ag , AgBr ELECTRODE ON THE MOLAR CONCENTRATION ( E p / V ) AND MOLE FRACTION (E$/V) SCALES CALCULATED AT 25 O C TETRAHYDROFURAN + WATER SOLVENT MIXTURES AT 5-45 OC AND THE STANDARD POTENTIALS OF 0 5 10 15 20 25 30 35 40 45 50 55 60 65 70 75 80 85 90 95 1 00 7.105 6.998 6.890 6.773 6.645 6.477 6.241 5.876 5.414 4.858 4.229 3.514 2.741 1.867 0.928 - 0.059 - 1.082 -2.131 -3.199 -4.296 -5.421 4.903 4.707 4.682 4.958 5.499 6.336 7.348 8.152 8.98 1 9.806 10.647 1 1.326 1 1.993 12.516 12.988 13.340 13.619 13.825 13.997 14.127 14.164 3.027 3.413 3.807 4.200 4.520 4.693 4.813 4.873 4.900 4.893 4.860 4.793 4.700 4.607 4.560 4.567 4.673 4.860 5.073 5.300 5.567 0.070 90 0.069 75 0.068 56 0.067 25 0.065 81 0.063 95 0.061 38 0.057 49 0.052 61 0.046 76 0.040 16 0.032 68 0.024 59 0.015 47 0.005 67 -0.004 62 -0.015 30 -0.026 26 -0.037 44 -0.048 93 -0.060 72 ~ 0.135 35 0.134 45 0.133 48 0.132 52 0.131 59 0.130 95 0.130 88 0.131 98 0.133 92 0.136 65 0.139 94 0.143 91 0.148 25 0.153 37 0.158 85 0.164 49 0.170 11 0.175 54 0.180 60 0.185 27 0.189 33 As in our previous following the procedure reported by Mussini et aL8 the standard e.m.f. in THF-rich solvents have been determined.STANDARD ELECTRODE POTENTIAL The least-squares values of the standard potential e (molality scale) of the Ag,AgBr electrode found in this investigation are summarized in table 1, along with the values for water as solvent.* The ion-size parameters that gave a satisfactory linear extrapolation were in the range 0.52 to 0.66 nm depending on both temperature and solvent composition. The standard deviation in is kO.05 and k0.09 mV for solvents containing 5-50 and 55-100% (w/w) THF, respectively. The values of e obtained for each solvent were fitted by the method of least-squares to = ~ - b ( t - 2 5 ) - c ( t - 2 5 ) ~ (1) where t is the temperature in O C .The parameters a, b and c are given in table 2 for each solvent, along with the values for water as the s01vent.~ Values of calculated by eqn (1) and the experimental values (table 1) generally agree within f0.11 V. The standard potentials @ on the concentration and AT$? on the mole fraction scales were computed at 25 OC with the help of the usual relation^,^ and are also included in table 2. Reported standard potentials of the Ag , AgBr electrode in THF + water solvent mixtures3 and our corresponding new values are collected in table 3 for comparison.TABLE COMPARISON BETWEEN THE NEW VALUES OF Erne AND PREVIOUSLY REPORTED VALUES BY ROY et aL3 temperaturePC THF W % ) 5 10 15 20 25 30 35 40 45 10 0.076 86 0.076 79 0.075 10 30 0.075 08 0.063 35 50 0.061 60 reported values3 0.057 98 0.075 11 0.073 23 0.071 13 0.068 88 0.066 44 0.063 76 0.060 94 this work 0.075 08 0.073 19 0.071 12 0.068 89 0.066 50 0.063 79 0.060 97 0.057 97 reported values3 0.072 30 0.069 34 0.066 16 0.062 41 0.058 52 0.054 50 0.050 19 0.045 53 this work 0.045 88 0.072 32 0.069 29 0.066 02 0.062 42 0.058 59 0.054 61 0.050 35 reported values3 0.058 23 0.053 08 0.047 99 0.042 30 0.036 69 0.030 98 0.025 31 0.019 34 thls work 0.019 15 0.057 07 0.052 44 0.047 55 0.042 29 0.036 83 0.031 14 0.025 251262 STANDARD POTENTIALS OF THE Ag,AgBr ELECTRODE The new values of are in good agreement with 22 out of 27 values obtained by Roy et aL3 for the 10, 30 and 50% THF+water mixtures at 5-45 OC.The differences range from 0.01 to 0.19 mV. However, Roy et aL3 expressed their values obtained in each solvent as a function of temperature, but the values calculated by their equations3 are not in complete agreement with their experimental values rep~rted.~ For example, in the 30% THF solution at 55 O C , the calculated value is 0.03507 V, whereas the experimentally reported value is 0.02930 V, and so there is a difference of 5.77 mV. Moreover, a minus sign should be added to the value of the parameter co (table 13) appearing in eqn (12) of their paper,3 and thus its value must be - 5.728 V K-2. STANDARD THERMODYNAMIC FUNCTIONS FOR THE CELL REACTION The standard thermodynamic functions for the cell reaction 4H2 (8, 1 atm)+AgBr (s) = Ag (s)+HBr (solvated) (2) were evaluated from the temperature variation of the standard molal e.m.f.in THF + water solvent mixtures. Thus, the standard changes of free energy (AG@) were evaluated using AG* = -nF@ = - F [ a - b ( t - 2 5 ) - - ~ ( ~ - 2 5 ) ~ ] . (3) The standard thermodynamic functions of the cell were computed at 5-45 OC by the usual relation~~9~ and these are recorded in table 4. The values of AGe are accurate to +9 J mol-l. It is evident from table 4 that the standard free energy changes for the cell reaction increase with an increase in either the THF content in the solvent mixture or the temperature of the solvent system. The standard enthalpy and entropy changes are all negative. At 25 O C , for example, the values of A P increase to a maximum at ca.10% THF and thereafter decrease, while the values of AH@ increase to a maximum at ca. 10 % THF, then decrease to a minimum at ca. 70 % THF, and thereafter increase again with increasing concentration of THF. STANDARD THERMODYNAMIC QUANTITIES FOR THE TRANSFER PROCESS The standard thermodynamic quantities for the transfer of 1 mole of HBr from the standard state in water to the standard states of the respective solvents HBr (in water) = HBr (in respective THF + water solvents) (4) were obtained from the temperature variation of standard e.m.f. of the cell on the mole fraction scale to eliminate the Gibbs energy change as a result of concentration changes in the transfer processlo EN8 = a’- h’ T - c’ T2. ( 5 ) The standard changes of Gibbs free energy (AGt@) can thus be represented as a function of temperature (in K) by (6) The least-squares values of the parameters of eqn (5) and (6) are given in table 5.The proper choice of a function to express the thermodynamic quantity as a function of temperature has been discussed in some detail by Ives and Marsden.ll The standard changes of enthalpy (AHt@), entropy (AS,@) and heat capacity (AC?) for the transfer process of HBr from water to the respective solvents were obtained by applying the F(wE$-sE$) = AGF = A - BT+ CT2.TABLE 4.-sTANDARD MOLAL THERMODYNAMIC FUNCTIONS OF THE CELL REACTION IN VARIOUS TETRAHYDROFURAN -k WATER SOLVENT MIXTURES AT 5-45 "C tempera- THF (wt %) ture 1°C 0 10 20 30 40 50 60 70 80 90 100 -AG*/J rnol-I 6768 5947 6043 506 1 5224 4080 4310 3006 3302 1838 -AHe/J rnol-I 25 610 29 304 28 287 31 959 31 059 34 708 33 926 37 551 36887 40488 -ASe/J K-l mol-l 67.7 84.0 77.2 93.3 86.7 102.7 96.1 112.1 105.6 121.5 5 15 25 35 45 7685 7299 6855 6353 5792 7404 7063 6648 61 59 5597 7298 6898 641 1 5837 5176 7254 6684 6022 5266 441 8 4778 3756 2645 1442 149 3226 2105 895 - 402 - 1787 1 404 225 - 1044 - 2403 - 3852 - 581 - 1785 - 3087 - 4486 - 5983 -2712 -3918 - 5230 - 665 1 -8178 5 15 25 35 45 17 593 19 247 20 960 22 730 24 560 15 883 17 963 20 116 22 343 24 644 17 204 19 673 22 230 24 874 27 606 21 807 24 437 27 159 29 975 32 883 31 918 34 486 37 145 39 894 42 734 33 187 35 678 38 258 40 925 43 681 32 937 35 490 38 134 40 867 43 691 31 537 34 309 37 178 40 146 43 212 29 324 32 366 35 515 38 772 42 136 5 15 25 35 45 35.6 41.5 47.3 53.1 59.0 30.5 37.8 45.2 52.5 59.9 35.6 44.3 53.1 61.8 70.5 52.3 61.6 70.9 80.2 89.5 97.6 106.6 11 5.7 124.8 133.9 107.7 116.5 125.3 134.1 142.9 113.4 122.4 131.4 140.4 149.4 115.5 125.3 135.0 144.8 154.6 115.2 125.9 136.7 147.4 158.11264 STANDARD POTENTIALS OF THE Ag,AgBr ELECTRODE TABLE 5.-vALUES OF THE CONSTANTS a’, b’ AND C’ OF EQN (5) FOR THE EVALUATION OF @ IN TETRAHYDROFURAN + WATER SOLVENT MIXTURES AT 5-45 O C AND THE VALUES OF THE CONSTANTS HBr FROM WATER TO TETRAHYDROFURAN + WATER MEDIA A, B AND C OF EQN (6) FOR THE EVALUATION OF THERMODYNAMIC QUANTITIES FOR TRANSFER OF THF - b’/ 10-4 c’/ 10-6 A/102 B/J K-’ C/10-2 (wt %) -d/10-’ V V K-’ V K-2 J mol-l mol-l J K-2 mol-l 0 10 20 30 40 50 60 70 80 90 100 5.1847 12.9923 17.1396 14.6353 1 1.3669 7.2292 3.2818 0.8837 2.0168 6.5626 12.6780 6.2245 11.2311 14.81 14 14.8691 13.9303 12.22 10 10.1416 8.5643 8.9034 1 1.2687 14.5001 3.027 3.807 4.520 4.813 4.900 4.860 4.700 4.560 4.673 5.073 5.567 - 75.3313 115.3455 91.1837 59.6485 19.7256 - 18.3607 - 4 1.4986 - 30.5653 13.2940 72.2984 - 48.3066 82.8506 83.4072 74.3491 57.8573 37.7940 22.5761 25.8476 48.6695 79.8468 - 7.5258 14.405 1 17.2321 18.0716 17.6856 16.1419 14.791 1 15.8814 19.7407 24.507 1 usual thermodynamic relations4 to eqn (6).The standard transfer thermodynamic quantities calculated at 5-45 O C are collected in table 6. The values of AGt* are accurate to f 17 J mol-l. The standard Gibbs free energy of transfer is an index of the differences in interactions of the ions (for example, H+ and Br-) and the solvent molecules in the two different media.The values of AGF decrease negatively, pass through minima (at ca. 30 and 20% THF at 5-25 and 35-45 OC, respectively) and thereafter increase to positive values with increasing THF content in the solvent. The observed negative decrease in AGte values suggests that the transfer of HBr from water to the THF + water solvents is increasingly favourable. The negative values of AGP support the view that water is less basic than the mixed solvent, whereas the positive AGte values indicate that HBr is in a higher free-energy state in pure THF than in water, and therefore the transfer process is not spontaneous. Hydrobromic acid thus is more strongly stabilized in water-rich solvents (with maxima at ca. 30 and 20% THF at 5-25 and 35-45 O C , respectively) than in water, whereas for THF-rich solvents the solute is more strongly stabilized by solvation with water molecules.The standard transfer enthalpy and entropy show similar trends. At 25 O C , for example, the values of AHte and ASP decrease, pass through minima at 70 and 90% THF, respectively, and thereafter increase with increasing THF content in the solvent. The values of AHP and ASte could provide an insight into the solvent structure. The transfer process of ions from water to a mixed solvent includes a number of changes connected with building up and breaking down the structure.12 Further, the structure- forming processes are exothermic and accompanied by a decrease in entropy, and the structure-breaking processes are endothermic and lead to an increase in entropy.The negative and decreasing values of AHte and ASte assume that ions are breaking the water structure more effectively than in the mixed solvent. Water alone is therefore a more structured solvent than the THF + water mixtures. On the other hand, the positive entropy and enthalpy of transfer of HBr from water to water-rich solvents can be attributed to a greater degree of structure breaking by HBr in these solvents than in water. The values of the heat capacity (ACF) are all negative and decrease with increasingTABLE 6.-sTANDARD THERMODYNAMIC QUANTITIES (MOLE-FRACTION SCALE) FOR THE TRANSFER OF HBr FROM WATER TO TETRAHYDROFURAN +WATER SOLVENT MIXTURES AT 5-45 O C tempera- THF (wt %) ture /"c 10 20 30 40 50 60 70 80 90 100 F F m r m m AG,e/J.rnol-l - 438 140 - 15 676 444 1245 937 1845 1467 2479 11 710 14 325 12 712 15 239 13 749 16 185 14 821 17 164 15 929 18 175 40.5 52.0 44.1 55.2 47.6 58.5 51.1 61.7 54.7 64.9 98.4 89.8 101.9 93.0 105.5 96.3 109.0 99.5 112.5 102.7 - AHte/J mol-l -AS,*/J K-l mol-l -ACF/J K-l mo1-l 5 15 25 35 45 -81 - 138 - 180 - 206 -218 - 365 - 378 - 362 -317 -244 - 749 - 608 -431 -221 25 - 734 - 454 - 138 214 603 1014 1626 2267 2938 3639 204 1 2682 3355 4059 4795 3065 3696 4367 5077 5827 398 1 4570 5209 5896 6633 5 15 25 35 45 - 171 1 - 1284 - 843 - 387 84 - 390 426 1271 2144 3046 4214 5190 6200 7245 8324 8017 9040 10 100 11 195 12 327 15 593 16 431 17 298 18 195 19 121 15 344 16 243 17 174 18 137 19 132 13 943 15 061 16 219 17 416 18 652 11 731 13 119 14 555 16 041 17 576 5 15 25 35 45 - 6.4 - 4.9 - 3.4 - 1.9 - 0.4 - 2.7 0.2 3.0 5.9 8.8 12.5 15.9 19.3 22.8 26.2 26.2 29.8 33.4 37.0 40.6 59.7 62.7 65.6 68.6 71.5 62.5 65.7 68.9 72.0 75.2 61.1 65.1 69.0 73.0 76.9 56.5 61.4 66.3 71.2 76.1 5 15 25 35 45 41.9 43.4 44.9 46.4 47.9 80.1 83.0 85.9 88.8 91.7 95.9 99.3 102.8 106.2 109.6 100.5 104.1 107.8 111.4 115.0 82.3 85.2 88.2 91.2 94.1 88.3 91.5 94.7 97.9 101.1 109.8 113.8 117.7 121.7 125.6 136.3 141.2 146.1 151.0 155.91266 STANDARD POTENTIALS OF THE Ag,AgBr ELECTRODE TABLE 7.-ELECTRICAL AND CHEMICAL PARTS OF THE THERMODYNAMIC QUANTITIES ACCOMPANYING DIFFERENCE BETWEEN THE FREE ENERGIES OF TRANSFER OF THE CHLORIDE AND BROMIDE IONS THE TRANSFER OF HBr FROM WATER TO TETRAHYDROFURAN + WATER SOLVENT MIXTURES AND THE (AGte’/J mol-l), ALL CALCULATED AT 25 OC 10 20 30 40 50 60 70 80 90 100 485 969 1404 1632 1915 2746 4773 8397 14 367 - 73 I66 300 486 746 1126 1694 2555 4052 7449 - 253 - 528 - 731 - 624 - 302 119 573 800 315 - 2240 47 169 339 599 843 1149 1290 1423 1294 804 - 890 1102 586 1 950 1 12 906 15 036 16 008 15 751 14 925 13 751 0.4 1.1 2.1 3.6 5.3 7.6 10.0 13.3 17.9 27.7 - 3.8 1.9 17.2 29.8 42.3 50.9 55.6 55.6 51.1 38.6 temperature.The values of AC? decrease, pass through minima at ca. 40% THF, then increase to maxima at ca. 70% THF, and thereafter decrease again with increasing THF content in the solvent. To study the ion-solvent interaction, the method adopted by Khoo and Chan13 was followed. In this method,14 consider a function AGte’ on the mole-fraction scale given (7) The difference between the free energies of transfer of hydrochloric2 and hydrobromic acids gives the difference between the free energies of transfer of the chloride and bromide ions.The values of Act*’ so calculated at 25 OC are given in table 7. The values of AGP’ are positive for all the solvents and increase with increasing THF content in the solvent. This is qualitatively in agreement with the Born theory, which predicts that the bromide ion should be in a lower free-energy state than the chloride ion in mixed solvents of lower dielectric constant than water.14 Therefore, the Born equation may be expected to show an increasingly better fit as the THF content of the solvent is increased. The Gibbs energy of transfer may be divided into an electrostatic part AGZ, caused by the change in the dielectric constant of the medium, and a chemical part AGg, due to the difference in solvation and other ion-solvent interactions:’.by AG,,’ = AGp(HC1) - AGP(HBr) = AGP(Cl-) - AG,e(Br-). AGF = AG: + AGg Similar equations exist for the other thermodynamic quantities AHP and ASt,. Both the electrostatic and chemical parts of the standard thermodynamic quantities for the transfer process were calculated by using the usual relations,lV4 and the values so computed at 25 OC are given in table 7. The values of AGg are all positive and increase with increasing THF content in the solvent. The chemical part of the free energy (which appears to be negative for water-rich solvents and pure THF) decreases, passes through a minimum at ca.30% THF, then increases to a maximum at around 70% THF and thereafter decreases with increasing THF concentration in the solvent, finally becoming negative in anhydrous THF. As AGg is an indicator of the acid-base properties of mixed solvents,lV4 theM. M. ELSEMONGY, I. M. KENAWY A N D A. S. FOUDA 1267 negative AG: values indicate that the chemical reaction in the transfer process is spontaneous, and the spontaneity (and so the basicity) of the water-rich solvents increases, reaches a maximum at cu 30% THF and thereafter decreases with increasing THF concentration in the solvent. The more negative value of AG: obtained for anhydrous THF indicates that the transfer process is more favourable. On the other hand, the positive values of AG: obtained for THF-rich solvents indicate that the transfer process is unfavourable.Thus, as far as chemical reaction or solvation is concerned, hydrobromic acid is in a lower Gibbs-energy state in water than in these TH F-ric h solvents. The electrostatic parts of the enthalpy and entropy of transfer are all negative, whereas their chemical contributions, which appear to be positive only for 10% THF solvent, decrease, pass through minima at ca. 70% THF and thereafter increase with increasing THF content in the solvent. The large negative AHg values for THF-rich solvents reflect the smaller enthalpy changes involved in creating a correct configura- tional change of the solvent for the transfer process. This view is further supported by the negative values of ASg, which are associated with structural changes as far as the chemical interaction or solvation @r the transfer process is concerned. This phenomenon produces overall order and hence AS2 values are negative. We thank Mrs Laila Abu Elela for assistance with the-computations and helpful suggestions. K. H. Khoo, J. Chem. Soc. A, 1971, 2932. M. M. Elsemongy, Electrochim. Acta, 1978, 23, 881. R. N. Roy, E. E. Swensson and G. LaCross, J. Chem. Thermodyn. 1975, 7 , 1015. M. M. Elsemongy, A. Fouda and M. F. Amira, Electrochim. Acta, 1981, 26, 255. M. M. Elsemongy and A. S. Fouda, J. Chem. Thermodyn., 1981, 13, in press. C. Carvajal, K. J. Tolle, J. Smid and M. Szwarc, J. Am. Chem. Soc., 1965, 87, 5548. F. E. Critchfield, J. A. Gibson and J. L. Hall, J. Am. Chem. Soc., 1953, 75, 6044. T. Mussini, C. M. Formaro and P. Andrigo, J. Electroanal. Chem., 1971, 33, 177. H. S. Harned and B. B. Owen, The Physical Chemistry of Electrolytic Solutions (Reinhold, New York, 3rd edn, 1958), p. 459. lo R. A. Robinson and R. H. Stokes, Electrolyte Solutions (Butterworths, London, 2nd edn, 1965), l1 D. J. G. Ives and P. D. Marsden, J. Chem. Soc., 1965, 649. l2 K. Bose, K. Das and K. K. Kundu, J. Chem. Soc., Faraday Trans. 1 , 1978, 74, 1051. l 3 K. H. Khoo and C. Chan, Aust. J. Chem., 1975,28, 721. l4 B. K. Das and P. K. Das, J. Chem. Soc., Faraday Trans. 1, 1978, 74, 22. p. 353. (PAPER 1/969)

ISSN:0300-9599    

DOI:10.1039/F19827801257

出版商:RSC    

年代:1982

数据来源: RSC

摘要:

J. Chem. SOC., Faraa'ay Trans. 1, 1982, 78, 1269-1278 Triplet State Electron Spin Resonance Studies of Aryl Cations Part 5.-The Aryl Cation/Aryl Radical Product Ratio in Photolysis of Arenediazonium Salts at 77 K1 BY HANNA B. AMBROZ AND TERENCE J. KEMP* Department of Chemistry and Molecular Sciences, University of Warwick, Coventry CV4 7AL Received 18th June, 1981 An e.s.r. study has been made of the factors (solute concentration, matrix, irradiation wavelength and time, post-photolysis storage) affecting the relative yields of 3Ar+ and Ar. in the photolysis of an arenediazonium salt at 77 K. The principal findings are: (i) (in contrast with the situation prevailing in solution) photoheterolysis is the sole primary process; (ii) production of Ar - during irradiation comes from photoreduction of 3Ar+ ; (iii) there is a very powerful solute-concentration dependence of the photo- conversion of 3Ar+ into Ar.(which does not proceed at very low solute concentrations); (iv) conversion of 3Ar+ into Are proceeds slowly in the dark. Mechanistic aspects of these observations are discussed. We have shown previously that the photolysis at 77 K of arenediazonium salts suitably substituted with good n-donors, eqn (l), yields two principal reactive intermediates, namely the triplet state aryl cation, 3Ar+, and the aryl radical, Are, both of which can be trapped and identified by their e.s.r. spectra,2 which are especially characteristic for 3Ar+. In the case of other arenediaz- onium salts, the electron configuration of Ar+ is singlet and no e.s.r. signal is expected or observed for this species, and only the absorption of Are is apparent.The ratio [3Ar+]/[Ar *I, denoted T/R, which could reflect the relative importance of eqn (1) and (2) if the trapping efficiencies of the intermediates are identical, has been measured2 and varies from 4.0 (for the 3-methoxy-4-morpholino compound) to much smaller ratios than 1 .O for most other arenediazonium salts. These figures refer to our normal experimental procedure of photolysis using irradiation from an unfiltered 200 W high-pressure Hg arc for periods of several hours using microcrystalline samples. We also noted during previous experiments that T / R varied with the irradiation time. When this point was explored in more detail, it was found that the variation could be extreme, and that T / R was also influenced strongly by (i) the wavelength of radiation, (ii) the type of matrix, (iii) the concentration of diazo salt and (iv) storage of photolysed sample in the dark for some hours at 77 K.In view of the current interest 1269c h) 4 0 TABLE 1 .-DEPENDENCE OF T / R RATIOS UPON EXPERIMENTAL CONDITIONS time of photolysis/min storage U.V. light visible light time in matrices dark/h 4 10 30 60 120 4 10 20 30 60 120 microcrystalline 0 powder 24 48 72 LiCl/H,O glass 0 (high [Dz+l) 5 24 48 72 (low [ Dz+]) 24 48 72 LiCl/H,O glass 0 1.21 0.442 0.304 0.273 0.036 0.015 0.014 0.013 9.17 3.87 2.67 2.51 - 0.736 0.333 0.232 0.172 0.051a 0.025 0.020 0.01 8 2.93 1.38 1.16 0.95 - 0.469 0.235 0.164 0.119 0.1 34a 0.071 0.059 0.053 0.686a 0.300 0.229 0.198 - 0.1 372a 0.149 0.107 0.079 0.014a 0.0067 0.0067 0.0060 0.3Sa 0.142 0.116 0.103 - 0.214a 0.118 0.07 1 0.052 0.036a 0.023 0.019 0.017 0.63tIa 0.38 1 0.309 0.303 - 3.41 4.40 1.68 2.11 1.70 1.87 1.41 1.68 0.674 3.64 0.366 1.96 0.301 1.84 0.306 1.53 0.275 1.37 3.88 3.50 1.91 1.73 1.58 1.69 1.93 1.73 8.02 9.73 3.60 3.99 2.38 2.63 2.19 2.42 - 1.42 - 0.585 - 0.522 - 0.485 - 3.50 - 1.54 1.38 - 1.28 - - - 7.18 5.16 2.6 1 2.23 1.59 1.73 1.43 1.56 0.083 0.041 0.045 0.027 0.041 0.023 0.038 0.022 9.50 8.58 4.04 3.24 3.67 2.94 3.38 2.37 - - a Different instrumental settings (1.0 mT field modulation at 10, 20 or 26 dB microwave power).H.B. AMBROZ A N D T. J. KEMP 1271 in solid-state photo~hemistry,~ we now report these rather unusual phenomena in detail, taking as our example the behaviour of 2,4,5-trimethoxybenzenediazonium hexafluorophosphate, Dz+PF;, which was typical in giving a very small value for T / R on prolonged U.V.photolysis. EXPERIMENTAL Samples were of three kinds, namely (i) microcrystalline, (ii) solutions in LiCl/water/acetone (ca. 6.0 mol dmP3 LiCl in 50/50 v/v solvent mixture) at high [Dz+PF;] (ca. 0.5 mol dm-3) and (iii) solutions in water/acetone/LiCl (ca. 6.0 mol dm-3) at low [Dz+PF;] (ca. 0.1 mol dm-3). Samples were made up and handled as in ref. (2a). PHOTOLYSIS TIMES Photolysis times were varied from a few minutes (i.e. much shorter than we have used before) to 2 h: the e.s.r. signals achieved during short photolyses were naturally of much weaker intensity and were measured with particular care (using slow scans, large time constants, etc.).All photolyses and measurements were conducted at 77 K, and e.s.r. spectra were recorded immediately afterwards (and then again after several set times of storage in the dark). Irradiation was performed using two types of light source: in one case this was a 100 W high-pressure Hg arc (unfiltered) and in the second a 200 W tungsten bulb (ie. giving no light at 2 < 350 nm); radiation from these sources is referred to as ‘u.v.’ and ‘visible’ light, respectively. E.S.R. MEASUREMENTS [3Ar+] was measured from the intensity of the sharp Hmin feature, while [Ar.] was measured from that of the integrated resonance near the free-spin value. The majority of experiments were conducted under the following conditions : field modulation, 2.0 mT ; frequency modulation, 100 kHz; time constant, 0.1 s; microwave power, 10 dB.RESULTS AND DISCUSSION Table 1 summarises the results of varying (i) irradiation wavelength, (ii) photolysis time, jiii) matrix and (iv) storage time of the sample following irradiation. Several trends are apparent which we discuss below. Fig. 1-3 illustrate the dramatic affect of using visible (as opposed to u.v.) photolysis for a microcrystalline sample. While U.V. light gives [Are] > [3Ar+] and a T / R ratio which decreases monotonically with photolysis time, visible light produces a large predominance of 3Ar+ and a T / R ratio which initially increases with photolysis time and then maximises before decreasing. EFFECTS OF IRRADIATION WAVELENGTH A N D TIME I N DIFFERENT MATRICES In all matrices, the use of U.V.light results in higher [Are] and lower T / R ratios than obtained with visible light under comparable conditions, with a single exception; when [Dz+PF;] < 0.1 mol dmW3 in a LiCl glass, [Are] becomes extremely small, i.e. T / R exceeds 20 (table 2). MICROCRYSTALLINE MATRIX Under U.V. irradiation, the T / R ratio is maximal (near unity) at shortest irradiation times, but falls rapidly as photolysis proceeds to an extent requiring adjustment of instrumental settings and some consequent inaccuracy in the measurement of T/R (probably ca. 50% when T / R z 0.1). Visible light produced lower concentrations of paramagnetic species in comparable periods, but the tendency to much higher T / R1272 TRIPLET STATE E.S.R.STUDY OF ARYL CATIONS TRIPLET RADICAL FIG. 1.-E.s.r. spectra of 3Ar+ and Ar. under different irradiation conditions: (a) 30min visible light photolysis, (6) 2 h U.V. photolysis. Field sweep, 0.7 T [spectrum (a) is displaced to right for clarity]. time of photolysis/h FIG. 2.-Dependence of T / R ratios (%) upon times of photolysis and storage in dark (typical for U.V. irradiation of microcrystalline samples).H. B. AMBROZ AND T. J. KEMP 1273 time of photolysis/h FIG. 3.-Dependence of T/R ratios (%) upon times of photolysis and storage in dark (typical for oisible light photolysis of microcrystalline samples). ratios was unequivocal, as was the appearance of a maximum after 30 min of nearly 10. These observations contrast strongly with those found with glassy solutions.LiCl GLASS (HIGH [Dz+]) When [Dz] x 0.5 mol dm-3, U.V. irradiation gives T / R ratios ’< 0.1, which cannot be measured sufficiently accurately to permit many conclusions to be drawn about the effect of photolysis time apart from the obvious one that 3Ar+ production is insignificant under these conditions. However, visible light yielded T/R ratios of up to a maximum of ca. 3.5 after 10 min irradiation. LiCl GLASS (LOW [Dz+]) When [Dz+] x 0.1 mol dm-3, T / R ratios exceeding 9 are obtained with both U.V. light (after short photolysis, but decreasing continually on longer irradiation) and visible light (maximising at 9.5 after 60 min photolysis). The striking contrast with the more concentrated glassy solutions of [Dz+] under U.V.irradiation stimulated a more detailed examination of the dependence of the T / R ratio on [Dz+] summarised in table 2. This (and similar experiments not detailed) reveal a remarkably abrupt change in T / R between 0.08 and 0.12 mol dm-3 [Dz+] covering nearly two orders of magnitude (making T/R extremely sensitive to slight variation in [Dz+]). Such an extreme changeover suggests a microenvironmental effect ; thus Dz+PF; is almost insoluble in LiCl/H,O, requiring acetone co-solvent for solubilisation, and at high [Dz+] the solute ions are probably in clusters solvated by acetone (especially after freezing) and ready electron transfer from the counter-ion to 3Ar+ can occur following1274 TRIPLET STATE E.S.R. STUDY OF ARYL CATIONS TABLE 2.-cONCENTRATION DEPENDENCE OF T / R RATIOS AFTER 4 min U.V.PHOTOLYSIS OF Dz+PF; IN LiCl GLASS (I, 11, I11 correspond to three separate series of measurements) [Dz+]/mol dm-3 (approximate) I I1 I11 comments 0.5 0.049 0.075 0.036 below 10% (i.e. too low to be significant) 0.25 0.169 - - - 0.125 - 0.1 - 3.24 9.17 1.27 more than 1000% ( i e . too high to be 0.063 26.90 17.91 0.03 1 41 S O - 0.015 27.40 - measured accurately) photolysis. At lower [Dz+], the solute exists as individual ion-pairs, which undergo photolysis to the separated ions, viz. 3Ar+ - * - N, - - PF;. The situation with visible light irradiation is less extreme : taking the maximum value of T/R realised, this shows the matrix-dependence : LiCl/H,O LiCl/H,O microcrystalline ([Dz+] 0.5 mol d ~ n - ~ ) < ([Dz+] 0.1 mol dm-3) - - 3.6 9.5 9.7 In summary, at [Dz+] > 0.5 mol dm-3, T/R < 0.1 and 3Ar+ is produced in insignificant quantity, while at [Dz+] < 0.8 mol dm-3, T / R > 20 and Ar is barely formed.(Indeed, the figures for [Are] include absorption from a species centred at g = 2.0 with a total width of 13 mT and features indicating a radical-pair with separation ca. 6.6 A, suggestive of a species [Are - - - PF,].) EFFECT OF STORAGE TIME In all irradiated samples (irrespective of wavelength or time of irradiation, the matrix or diazo concentration), the ratio T / R decreases slowly on storage of the sample in the dark at 77 K. Averaging over all of the 31 individual samples studied, the value of T/R (taken as 1.0 at the initial measurements) changed as indicated in fig.4 for the three different types of sample studied. Clearly T / R declines rapidly over the first 24 h period of storage, but subsequently much more slowly. In one experiment (see table 1) with a sample in glassy LiCl of Dz+PF, (0.5 mol dm-3), we also measured T / R after only 5 h of storage, and found that much of the change in T/R occurred in this initial period. The universal decrease in T / R is to be associated with growth in [Ar-] on storage, coupled with decrease in [3Ar+]. We observed both of these changes many times, and fig. 5 is quite typical, although they were harder to quantify exactly than simple measurement of the ratio T / R . However, the following trends were discerned. (i) The largest relative growths in [Ar *] were found after visible-light irradiation for short periods, reaching nearly 3 times the original concentration achieved after photolysis.The percentage increase in [Are] was smaller for samples subjected to longer visible irradiation, especially for the glassy samples at low [Dz+PF;] when it was as little as 20%.H. B. AMBROZ AND T. J. KEMP 1275 0 24 48 72 96 storage time in dark/ FIG. 4.-T/R ratios after storage in dark in different matrices (initial T / R measured after photolysis taken as 1 .OO). Matrices: 0, microcrystalline powder; x , LiCl/water glass, ca. 0.1 mol dm-3 Dz+; A, LiCl/water glass, ca. 0.5 mol dm-3 Dz+. t -/- I I I I / / I J I 2 4 6 24 tstlh FIG. 5.-Storage-time profiles of [3Ar+] and [Arm] after photolysis: both [3Ar+] and 1Ar.J at storage time zero are taken as 1.0.T = 77 K. Sample: 0.08 M Dz+ in LiCl/H,O/acetone glass. (ii) Following U.V. photolysis, storage produced only a 30-40% increase in [Ar-1, and then only following short irradiation times; samples irradiated for 1-2 h produced only a 3-1074 increase in [Ar.] on storage. (It should be noted, of course, that if a sample has a very small T / R value on conclusion of photolysis, even complete conversion of 3Ar+ into Ar , the presumed sole mechanism for producing extra Ar , will increase [Ar = ] only marginally.) (iii) One special feature of the dilute solutions of Dz+PF; in glassy solution is that no conversion of 3Ar+ into Ar* proceeds in the dark. (iv) In some cases a small net decrease in [Ar J was found. This was typically of1276 TRIPLET STATE EAR. STUDY OF ARYL CATIONS the order of lo%, but the result seems genuine in that the decrease was reproducible and was large compared with signal drift.If there is some real concomitant decay of Are at 77 K, then this means that our estimates of the conversion 3Ar+ + Ar may be slightly lower than the true extent. To explore this point further, we investigated the photolysis of 2,4,6-trimethoxybenzenediazonium hexafluorophosphate, which we presume, from the established behaviour2b in aqueous solution at room temperatures, to give a (diamagnetic) singlet state Ar+ together with Are on photolysis (and, as before, no triplet Ar+ was observed). Storage of this material at 77 K following U.V. photolysis (10 min), both as microcrystalline samples and in glassy LiCl/H,O matrices, gave no enhancement of the Ar 0 signal, implying no conversion of singlet Ar+ into Are, and instead Ar= decayed ca. 10% over 72 h.We believe therefore that Are does decay on this time scale at 77 K. MECHANISMS FOR THE LIGHT AND DARK REACTIONS The most complete model takes the form: primary photochemical processes ArNiX- 14; 3Ar+ + N, + XL ArNzX-: Ar- +N,+X* secondary photochemical processes 3ArS $ ArX ( 5 ) Ar* 2 (Are)* 5 diamagnetic products (6) dark reactions 3Ar++ES +AR* +ES+* (7) (8) (9) ' 3Ar+ + X- + ArX Ar- + R - + diamagnetic products where ES is an electron source (the nature of which is discussed later) while ES+ is its electron-deficient counterpart. REACTIONS DURING PHOTOLYSIS The relative contributions of eqn (1)-(4) change as photolysis proceeds.Most significantly, and unexpectedly, no clear evidence for the primary photodissociation eqn (2) has been found for the arenediazonium salt examined; the ready production of Ar* during photolysis arises purely from the secondary reactions (3) and (4). The contributions of reactions (5) and (6) cannot be assessed accurately, but they are certainly significant and must influence the apparent rates of processes (I), (3) and (4)H. B. AMBROZ AND T. J. KEMP 1277 by providing competing routes for removal of 3Ar+. Under all the various experimental conditions utilised, i.e. different irradiation wavelengths, matrices and [Dz+], the photolysis features three principal stages: Stage I, i.e. the early stage when only process (1) is significant, Ar- is not detected and T/R assumes very high values.Stage 11, i.e. the middle stage when reactions (3) and/or (4) become significant, and Arm appears near g = 2.0. T / R falls steadily to reach 1.0, i.e. [3Ar+] = [Ar.]. Stage 111, i.e. the final stage. Now the secondary processes dominate, and a decrease in T / R to < 1.0 is observed. The exact point during photolysis at which a system reaches the various stages depends on such factors as the irradiation wavelength, the rigidity of the matrix, the energetics of the transformation 3Ar+ + Are and the (solid-state) oxidation potential of ES. The absence of any production of Are at the very lowest [Dz+] in LiCl/H,O/acetone glass even after prolonged irradiation suggests that a critical local concentration of 3Ar+ (and its counter-ion) is necessary for the 3Ar+ + Ar* conversion to proceed.(Cl-, despite its overwhelming preponderance, seems not to act as the major electron source: the function of PF; as the electron source is unambiguous for microcrystalline samples.) We associate the absence of 3Ar+ + Ar conversion at low [Dz+] with the separation of 3Ar+ and PF; by the N, molecule following photolysis of the ion-pair, Dz+PF;. At high [Dz+], the frozen solution contains clusters of the various solute ions, e.g. ArNzPF; 3Ar+ - - - N, - - - PF; hv -+ PF;+N,Ar and ready electron transfer between adjacent ions can cover. Stages 1-111 are more ‘compressed’ in time under U.V. irradiation: the rate of initial production of 3Ar+ is faster, as is the conversion (3) or (4). We expanded the sequence in time simply by reducing the incident U.V.light intensity by moving the sample either away from the source or nearer than its focus. Both procedures led to much higher T / R ratios, e.g. up to 1 .O for LiCl/H,O glass at high [Dz+], while the microcrystalline samples yielded T / R values of 5.5k0.2 (three determinations) for short (4 min) irradiations and 1.8 for a 10 min irradiation. To summarise, low T / R ratios are produced by the following conditions: (i) the use of a softer matrix, i.e. acetone/ H,O/LiCl (if [Dz+] is sufficiently high), (ii) U.V. light, (iii) higher light intensities and (iv) long photolysis times. A perplexing feature is the appearance of maxima in the dependence of T / R upon photolysis time using visible-light irradiation (irrespective of the matrix or diazo salt concentration), which cannot be understood solely in terms of eqn (1)-(4).These were found to be associated with a very strange build-up profile of [Are] with time; these profiles (in all matrices) display an initial build-up followed first by a well-defined, short decay region, and secondly by a slow further build-up. Individual growth profiles of [3Ar+] are not marked by any complications of this type, i.e. the maxima result solely from the kinetic behaviour of Are. One possible explanation involves two parallel reactions (3) or (4) involving ES sites of markedly different efficiency : the more efficient results in an initial fast build-up of Ar*, which subsequently disappears in a photostimulated decay, the less efficient gives a slower parallel build-up of Are and the sum of the two gives the ‘switchback’ profile. PF; - - - N, - - - 3Ar+1278 TRIPLET STATE E.S.R.STUDY OF ARYL CATIONS REACTION IN DARKNESS During post-photolysis storage in the dark, 3Ar+ slowly converts into Ar* (e.g. fig. 5) at rates (fig. 4) which are comparable for all samples irrespective of their varied photolysis conditions. While eqn (7) is clearly very important under appropriate reaction conditions, there are occasions when 3Ar+ disappears without a compensating increase in [Are], indicating the role of eqn (8) and (9): eqn (9) becomes more important at elevated temperatures but we detected slow decay in Ar* even at 77 K. Very few comparable studies exist, but in addition to 'normal' reactions of carbenes in matrices (to give dimers or cycloadd~cts),~* ti it has been reported recentlys that the quintet state formed by interaction of two benzoylphenylmethylene species decays on warming to 90 K to a new triplet state, regarded as the diradical remaining on combination of the two carbene species.Some comparisons should be made with the situation prevailing in solution thermolysis7 and photolysis8 of arenediazonium salts. Two principal types of product are obtained on either thermolysis7 or photolysis8 in MeOH, namely ArOMe (from Ar+) and ArH (from Are). The ratio ArOMe/ArH increases with (i) the presence of substituent electron (a and IC) (ii) the presence of 0,, which is expected to suppress radical-chain reactions,' (iii) the use of short-wavelength light (3 13 nm)* and (iv) the exclusion of added electron donors such as pyrene (which induce large yields of ArH).8 There is a contrast between the solution and solid-state photolyses in that photolysis at any wavelength promotes formation of 3Ar+ from solid samples +bile short-wavelength photolysis of methanol solution is required to produce ArOMe (longer wavelengths yielding ArH). We thank the S.E.R.C. for provision of the e.s.r. spectrometer, the gaussmeter and the cryogenics, and for support of H.B.A. Part 4. H. B. Ambroz and T. J. Kemp, J. Chem. SOC., Faraday Trans. 1, 1-982, 78, 725. * (a) H. B. Ambroz and T. J. Kemp, J. Chem. SOC., Perkin Trans. 2, 1979, 1420 and 1980, 768; (b) H. B. Ambroz and T. J. Kemp, Chem. SOC. Rev., 1979, 8, 353. ' (a) G. M. J. Schmidt et al., Solid State Photochemistry (Verlag Chemie, Weinheim, 1976); (b) J. Burdett and J. J. Turner, in Cryochemistry, ed. M. Moskovits and G. A. Ozin (Wiley, New York, 1976), chap. 1 1 ; (c) P. D. Fleischauer, in Concepts in Inorganic Photochemistry, ed. A. W. Adamson and P. D. Fleischauer (Wiley, New York, 1975), chap. 9; (d) G. M. Parkinson, M. J. Goringe, S. Ramdas, J. 0. Williams and J. M. Thomas, J . Chem. SOC., Chem. Commun., 1978, 134 and references cited therein. V. P. Senthilnathan and M. P. Platz, J. Am. Chem. SOC., 1980, 102, 7637. C-T. Lin and P. P. Gaspar, Tetrahedron Lett., 1980, 21, 3553. H. Murai, M. Torres and 0. P. Strausz, J. Am. Chem. SOC., 1980, 102, 5104. H. G. 0. Becker, G. Hoffmann and G. Israel, J. Prakt. Chem., 1977, 319, 1021. ' T. J. Broxton, J. F. Bunnett and C. H. Paik, Chem. Commun., 1970, 1363. (PAPER 1/994)

ISSN:0300-9599    

DOI:10.1039/F19827801269

出版商:RSC    

年代:1982

数据来源: RSC