摘要:

J . Chem. SOC., Faraday Trans. I , 1986,82, 89-102 Arrhenius Parameters for the Reaction HO, + C,H, - C,H, + H,O, over the Temperature Range 400-500 "C Roy R. Baldwin, Christopher E. Dean, (the late) Malcolm R. Honeyman and Raymond W. Walker" Chemistry Department, Uniuersity of Hull, Hull, N . Humberside HU6 7RX Studies have been made of the addition of C2H6 to mixtures of tetramethyl- butane (TMB)+O, over the temperature range 400-520 "C in both boric acid-coated and KC1-coated Pyrex vessels. From measurements of the relative yields of i-C,H, and C,H,, formed from TMB and C,H,, respectively, values of klo/k$ have been determined: HO, + C,H6 + H,O, + C2H, (10) HO,+HO, + H,O,+O,. (7) By use of low total pressures (ca. 15 Torr),? the contribution of OH attack to the overall formation of C,H, has been reduced to a very low level.At temperatures above 470 "C, the values of k,,/k; in KC1-coated vessels were noticeably above those obtained from the boric acid-coated vessels, and this has been attributed to possible gas-phase regeneration of radicals when H20, is removed at the KCl surface. Values of E,, -+E, = 82.7 f 1.5 kJ mol-l and Alo/Ai = 105.49+0.11 (dm3 mol-1 S - I ) ~ were obtained. Low-temperature values of k , have been partially re-interpreted and combined with Troe's single temperature value at 1100 K to obtain values of k , for use under the conditions of the present study. In this way, Arrhenius parametersofE,, = 85.6f2.5 kJ mol-landAl, = 1010.23*0.16 dm3 mol-l s-' were obtained. No other data are available for this reaction, but the parameters are in good agreement with A = 1010.38 dm3 mol-l s-' and E = 84.5f 8 kJ mol-l for the reaction of HO, radicals with TMB, both alkanes containing only primary C-H bonds.It has been common practice when developing chemical mechanisms for modelling hydrocarbon combustion phenomena in the region 300-1000 "C to assume that HO, radicals are effectively inert. Indeed, in the Shell model for auto-ignition the reaction R + 0, -+ conjugate alkene + HO, where R is an alkyl radical, is considered to play a key role as the main linear termination step.l More recently, by the use of sensitivity analysis with large comprehensive mechanisms, it has been demonstrated that H-atom abstraction from alkanes by HO, radicals, even though it is endothermic, is important in determining auto-ignition behaviour.For example, Westbrook and coworkers2 confirmed the importance of HO,+C,H,, in the oxidation of butane by use of a mechanism consisting of 238 elementary reactions amongst 47 species, and validated the mechanism by comparison of computer and experimental results from shock tubes, turbulent flow reactors, and premixed laminar flames. Similarly, use of a mechanism containing about 70 elementary reactions has provided clear evidence of the importance of the reaction HO, + CH, + CH, + H,O, in determining ignition delays3 in CH, + 0, mixtures, even t 1 Torr z 133.3 Pa. 8990 Arrhenius Parameters for HO, + C,H, though the reaction is ca. 60 kJ mol-1 endothermic. With alkanes containing tertiary C-H bonds, reactions of HO, radicals will play an even greater part in propagating combustion, and at temperatures above 400 "C will lead to degenerate branching through the homolysis of the hydrogen peroxide formed: H,O, + M + 20H + M.Unfortunately, very little information is available on HO, reactions with alkane^.^ No direct studies involving the measurement of HO, concentration have been made, and the limited number of data have all been obtained from relatively complex systems in the form of k(H0, + RH)/ki(HO, + HO,). From studies of the decomposition of tetra- methylbutane (TMB) in the presence of 0, in both KCl-coated and aged boric acid- coated vessels, Baldwin and coworkers5 found A = I .95 x 1 O 1 O dm, mol-' s-' and E = 81.7 8 kJ mol-1 over the temperature range 380-540 "C for HO, + TMB.Similar studies have given k(HO2+2,2,3-trirnethylbutane) = (3.1 kO.6) x lo5 dm3 mol-1 s-l at 500 oC.6 From studies of the addition of an alkane to HCHO + 0, mixtures, preliminary values of k(HO,+C,H,) = 1.6 x lo4, k(HO,+C,H,) = 5.1 x lo4 and k(HO,+i-C,H,,) = 9.0 x lo4 dm, mol-l s-l at 440 "C have been obtained.' From a photo-oxidation study, Alcock and Miles give k(H0, + 2,3-dimethylbutane) = 1.45 x 1 0, dm3 molF1 s-l at 100 "C. All these values are based on k(H0, + HO,) = 2.0 x los dm3 mol-1 s-l, indepen- dent of temperature. Lastly, studiess of the direct oxidation of a series of aldehydes have provided rate constants for HO, + RCHO at ca. 450 OC, both for attack at the aldehydic C-H position and at C-H sites in the R group. Recently, the decomposition of TMB in the presence of 0, has been developed5 as a reliable and controllable source of HO, radicals over the temperature range 38&520 "C.Detailed studies in both KC1-coated and aged boric acid-coated vessels have shown that the decomposition proceeds by a simple mechanism: TMB-2 t-C,H, (I)? (2) t-C,Hs +O, - i-C,H, + HO, (3) (4) ( 5 ) surface HO, - 4H,O + go2 HO, + TMB - H,O, + (CH,),CC(CH,),CH, (CH,),CC(CH,),CH, - t-C,H, + i-C4Hs 2H0, - H,O, + 0, (7) surface H,O, - H,O++O, OH +TMB - H,O + (CH,),CC(CH,),CH,. (9) In KC1-coated vessels, HO, and H,O, are efficiently destroyed at the surface so that the stationary concentration of H,O, is very low and is reached early in the reaction. Effectively, there is no autocatalysis due to reaction (8). By the use of relatively low pressures of TMB (d 4 Torr) and low total pressures (d 60 Torr), contributions to the removal of TMB by radical attack are limited to about 25% and reaction (9) is only important above 500 "C.However, the simultaneous presence of an efficient surface termination process and a second-order gas-phase termination process together with the occurrence of secondary initiation, makes the precise calculation of [HO,] both difficult and tedious.1° t The reaction numbers used are consistent with previous publication^.^R. R. Baldwin, C. E. Dean, M . R. Honeyman and R. W. Walker 91 With aged boric acid-coated vessels, surface destruction of HO, and H,O, does not occur so that k , = k6 = 0. The reaction is autocatalytic owing to the development of a non-stationary concentration of H,O, which slowly decomposes to give OH radicals.Although the [OH]/[HO,] ratio is higher in boric acid-coated vessels than in the KCl vessels, the highly reproducible nature of the surface, the occurrence of only gas-phase termination, and the fact that the rate constants for OH +alkanes are accurately known permits the use of the system for a study of HO, reactions, particularly if low total pressures (cu. 15 Torr) are used., This paper discusses the results obtained when C,H, is added to TMB+O, mixtures between 400 and 520 "C. Arrhenius parameters for reaction (10) are obtained for the first time. Fortunately, 99% of t-C,H,ll radicals and of C,H,12 radicals with 0, [reactions (2) and (1 l)] to give 1-C,H, and C,H,, respectively, under the conditions used, so that the production of OH radicals in alternative reactions is negligible.The basis of the method can be illustrated by considering the simplest possible mechanism involving reactions (l), (2), (7), (10) and (1 1) HO, + C,H, + H,O, + C,H, C,H, + 0, + C,H, + €30, from which the relative rates of formation of C,H, and 1-C,H, are given by d [ C , H ,] / d [ i -C, H ,] = k [ C ,H ,] / ( k k , )$[ TM B]i. The experimental results at short reaction times are consistent with eqn (i), but for more precise interpretation a computer treatment is required to allow for the presence of OH radicals. Experimental Details of the experimental procedure are given e1sewhere.j. l1 Reactions were carried out in cylindrical Pyrex vessels, 20f 1 cm in length and 5.1 f 0.1 cm in diameter.KC1 coatings were renewed at least once per week to prevent loss of efficiency for the destruction of peroxy species. Aged boric acid-coated vessels were prepared by carrying out repeated runs with H, + 0, mixtures at 500 "C, as described previ0us1y.l~ The yields of C,H, and of i-C,H, were measured by gas chromatography using a Perkin-Elmer Sigma 100 system, which incorporated microprocessor control over the important operating conditions. Electromagnetic valves coupled to a microprocessor unit (response time < 0.1 s) were used to admit the gases into the reaction vessel from the premixing bulb and to sample the products after a predetermined time interval. To avoid complications caused by secondary processes when determining rate constants, consumption of both TMB and C,H6 was never allowed to exceed 5 % .No systematic analyses were made for the minor products, which were investigated thoroughly earlier,l1- l2 but the possibility of unexpected results was covered by careful inspection of the chromatograms. Results Boric Acid-coated Vessels The yields of C,H, and of 1-C,H, over the first few percent of reaction were measured for a range of mixture composition at 400, 440, 470 and 500 "C. In the main phase of the study, the total pressure was maintained at the low value of 15 Torr to minimise the decomposition of H,O, and therefore the consumption of C,H6 by OH attack. The yields of both 1-C,H, and C,H, (more markedly) increase autocatalytically with time as shown in fig. 1 for several mixtures at 440 "C.Fig. 2 shows plots of [C,H,]/[i-C,H,] against time in the early stages of reaction for several mixtures at 440 "C. As expected from fig. 1, the values of the ratio increase steadily with time. Variation of the pressure of 0, between92 Arrhenius Parameters for HO, + C,H6 n &? v c .- Y 0.6 I I I I 1 0.4 0 - 2 0 2 1 0 25 50 75 100 ti s Fig. 1. Yields of i-C,H, and C,H, in an aged boric acid-coated vessel at 440 "C. 0, TMB = 4, C,H, = 5 , 0, = 6 Torr; x , TMB = 1, C,H, = 5, 0, = 9 Torr; 0, TMB = 4, C,H, = 1 , 0, = 10 Torr. 3-10 Torr has no noticeable effect on the ratio. As expected from the simple eqn (i), the product ratio is linearly dependent on the C,H, pressure and decreases markedly as the pressure of TMB is increased. No effect of variation in the vessel diameter between 2.0 and 5.1 mm is observed.Similar results are observed at all temperatures used, and fuller details are given later when discussing the interpretation of the data. A computer program was used for the evaluation of k,, from the [C,H,]/[i-C,H,] ratios. As previ~usly,~ differential equations are written for reactants and products, and stationary state equations are written for the radicals. It is convenient to express [HO,] in terms of the parameter G = k$[HO,J, since it is the ratios R, = k,/k$ and R,, = k,,/k$ that are directly involved in the computer treatment. Similarly, k, enters as the ratio R, = k,/k$, although for the boric acid-coated vessel R, = 0. In the mechanism, allowance is made for the 1 % of ethylene oxide and isobutene oxide formed from reactions (1 3) and (1 4), respectively : C,H, + 0, -+ C,H,O +OH (13) t-C,Hg + 0, + C,H,O +OH.(14) G is calculated from the stationary state equation for HO, and the differential equations are solved by the Kutta-Runga numerical integration method. Operation of the computer program requires values of the parameters k,, k14/k2, k13/'k11, R,, Rlo, kl,/kg and k,. The yields of 1-C,H, and C,H, are calculated for a range of mixture compositions and compared with the experimental values at selected times. Reactions (13) and (14) play only a minor part in the mechanism, and a value of k14/k2 = k13/kll = 0.010 hasR. R. Baldwin, C . E. Dean, M . R. Honeyman and R. W. Walker 93 0.55 0.45 - m 5 r;' C 0.35 ii" -. - P 3 - v 0.25 0.15 1 I I 1 ,x' 0 0 x / x 0 .--::.:: Y 0 20 L O 60 80 100 t l s Fig. 2. Variation of [C,H,]/[i-C,H,] with time of reaction in a boric acid-coated vessel at 440 "C. 0, TMB = 4, C,H, = 5 , 0 , = 6 Torr; x , TMB = I , C,H, = 5 , 0 , = 9 Torr; ( x lo), TMB = 4, C,H, = 1 , 0, = 10 Torr; A, TMB = 2, C,H, = 5 , 0, = 8 Torr. been taken from previous studies of the ethylene oxidel, and isobutene oxidell yields. The value of k, is known accurately for M = 0,, N, and H,, but the value for M = C,H, and TMB is a little more uncertain. Because of the relatively high pressures of TMB and C,H, in the mixtures used, TMB and C,H, will make a relatively high contribution to M , [the value of M in reaction (8)] which is given for any mixture by M , = 0.35 [O,] + a [TMB] + b [C,H,] (ii) where a and b are the efficiencies of TMB and C,H, relative to that of H,.Previous studies5 of the decomposition of TMB in the presence of 0, showed that optimum consistency of results in boric acid-coated and KC1-coated vessels was achieved with a = 2.5. These studies also gave accurate values for k , and R,. From studies15 of the separate addition of C,H, and TMB to slowly reacting mixtures of H,+O,, k,,/k, = 0.75 kO.1 at 480 "C, and because El, and E9 are very small and approximately equal, then the value will be effectively unchanged between 400 and 500 "C. Apart from the value of b in eqn (ii), which is discussed later, the only unknown parameter is Rlo. A range of values has been selected at each temperature and the relative yield of C,H, and 1-C,H, has been calculated at three times in the early stages of reaction and compared with the experimental values for each of the mixtures used.From an94 10 8 6 4 h F3.2 v m b 8 6 4 2 0 Arrhenius Parameters j o r HO, + C,H, \ 0.7 7 0.82 0 I ? X \x X z 0.40 0.45 0.50 I I I 0 5 3 -5 h * v C .- Y ca > -0 .- 2 E 5 0 - 5 0.20 0 -25 0.30 0.35 0.115 0.120 0,125 ( k I o / k , f ) / ( d r n 3 mol-' s-l)+ Fig. 3. Mean and r.m.s. deviations between observed and calculated results for boric acid-coated vessels: 0, r.m.s. deviation; x , mean deviation; (a) 400, (6) 440, (c) 470 and (4 500 "C. inspection of the r.m.s. deviation and mean deviation between experimental and calculated values of [C,H,]/[i-C,H,], the optimum value of R,, was obtained at each temperature. Although both deviations gave effectively the same values of R,, because of the rather flat minimum obtained from the r.m.s. deviations, the value of R,, giving zero mean deviation was considered the more reliable.Fig. 3 gives the variation of r.m.s. deviation with the value of R,, at the four temperatures used, when the other parameters are set at their optimum values, and shows that there is near coincidence in R,, between zero mean deviation and minimum r.m.s. values. Table 1 shows the calculated and experimental values of the ratio [C,H,]/[i-C,H,] at three times for the four mixtures used at each temperature when the optimum parameters are used. Table 1 also shows the optimum values of R,, as a function of temperature. A sensitivity analysis was carried out on the optimum values of R,,, the most important rate constant parameters in this context being k1,/k9 and the values of a and b in eqnR .R. Baldwin, C . E. Dean, M . R. Honeyman and R. W. Walker 95 Table 1. Calculated and experimental values of [C,H,]/[i-C,HJ mixture/Torr T/"C time/s TMB C,H, 0 2 400 440 470 500 900 900 900 300 300 300 90 90 90 90 l i } 15 l i } 15 l i ) 15 1;) 15 2 4 4 1 2 4 4 1 2 4 4 1 2 4 4 1 5 8 5 6 1 10 5 9 5 8 5 6 1 10 5 9 5 8 5 6 1 10 5 9 5 8 5 6 1 10 5 9 0.41 1 0.435 0.460 0.221 0.230 { 0.239 0.05 1 I 0.73 1 0.784 0.835 0.304 0.183 0.039 { 0.044 0.048 0.628 0.043 0.465 0.53 1 0.189 0.209 0.229 0.1 18 0.131 0.025( 1) 0.330 0.417 0.118 0.436 0.455 0.242 0.255 0.265 0.047 0.049 0.052 0.70 I 0.734 0.766 r.m.s. dev. = 7.2% 0.300 0.263 0.347 0.389 0.186 - 0.218 0.243 0.037 0.043 - 0.048 0.468 0.539 0.605 r.m.s.dev. = 3.5% 0.257 0.473 0.297 0.334 0.168 0.195 0.220 0.034 - 0.039 0.044 0.386 - 0.443 0.498 r.m.s. dev. = 4.3% 0.182 0.8 10 0.205 0.226 0.123 0.140 0.155 0.025(0) - 0.03 1 (2) - 0.267 0.296 0.325 r.m.s. dev. = 4.7% - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - 0.028( 5 ) - - - 4 F A R 196 Arrhenius Parameters for HO, + C,H, Table 2. Summary of sensitivity analysis" __ ~ ____ parameters changed min. ~ - r.m.s. change in change in b T/"C dev. E,,-+E, log,, (Al,/A$) k l , / k j k l , / k i (7:) ElO-iEi (%) 0.75b 0.50 1 .oo 0.75 0.75 1.5b 500 4.71 470 4.3 } 82.7rt 1.5 440 3.5 400 7.3 I 1.5 500 3.5 1 470 4'8 } 84.8+ 1.8 440 5.1 400 6.8 J 1.5 500 8.1 1 470 7'8 } 80.8k2.1 440 5.7 400 8.0 ' 470 440 "'} 3.7 84.05 1.6 400 7.4 0.9 500 2.1 500 5.1 1 470 4'4 } 81.2k1.6 440 4.2 400 7.3 ' c 0.810 0.473 5.49k0.11 { o.263 10.1 18 0.925 10.126 I 0.723 L0.113 0.881 0.519 0.121 0.753 l0.115 - 2.3 + 14.2 + 17.4 + 10.7 - 10.7 - 14.5 - 8.5 - 4.0 + 8.8 + 2.9 1.8 ~ a Energy in kJ mold', k, A in dm3 mol-l s-l.Recommended values for k l , / k , and b. (ii), which together control the amount of C,H6 removed by OH attack. Table 2 summarises the effect of relatively large changes in these parameters on the optimum values of Rlo. The effect on the rate constant values is relatively small at all temperatures, but the effect on the values of the Arrhenius parameters for reaction (10) is more significant because of the relatively small temperature range. As k12/kg = 0.75 f 0.10 from the addition studies15 and b almost certainly lies within the range 1.0-2.0, then values of El,-& = 82.7 f 1.5 kJ mol-I and loglo(Alo/A~) = 5.49 +O.1 1 (dm3 mol-1 s-l)$ are recommended. KCI-coated Vessels Studies in KC1-coated vessels, where the importance of OH attack on C,H, should be reduced considerably owing to the efficient destruction of HO, and H,O, at the surface, were carried out over the range 420-520 "C. Fig. 4 shows plots of [C,H,] and of [i-C,H,] against time together with plots of the ratio [C,H,]/[i-C,H,]. The plots do not show evidence of the autocatalysis observed with the boric acid-coated vessels, which is consistent with the computer prediction that in KC1-coated vessels the stationary concentration of H,O, is very low and is reached very early in the reaction.The same computer program was used to analyse the results and to determine the optimum value of Rlo. The significant difference in the mechanism is that the surface destruction of both HO, radicals and H,O, must be considered. Unfortunately, calculation of the values of k, and k6 is complicated by the interaction between the uniform profile of chain centres throughout the vessel, resulting from the homogeneous termination processes, and the diffusion profile created by an efficient surface termination.R. R. Baldwin, C. E. Dean, M . R. Honeyman and R. W. Walker 97 0.20 0.06 0.03 0 5 10 tl S 15 Fig. 4. Yields of 1-C,H, and C,H,, and values of [C,H,]/[i-C,H,] in a KC1-coated vessel at 440 "C. Upper: x , TMB = 5, C,H, = 2, 0, = 30, N, = 23 Torr; 0, TMB = 2, C,H, = 2, 0, = 30, N, = 26 Torr; A, TMB = 0.5, C,H, = 2, 0, = 30, N, = 27.5 Torr; 0, TMB = 2, C,H, = 10, 0, = 30, N, = 18 Torr.Lower: TMB = 2, C,H, = 10, 0, = 30, N, = 18 Torr; 0, 1-C,H,; x , C,H, ( x 3). TMB = 0.5, C,H, = 2 , 0 , = 30, N, = 27.5 Torr; A, i-C,H,; 0, C,H, ( x 10). These interactions can introduce errors of up to 25% in the HO, concentration. The calculation 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), (7) and (I 0) from HO, radicals which are not at a uniform concentration throughout the vessel, and since H202 is predominantly destroyed at the surface by a moderately efficient process, the total initiation process is not uniform.The procedure used to allow for this complication has been discussed el~ewhere.~~ lo Table 3 compares the results obtained for R,, with those already given for the aged boric acid-coated vessels. At 440 "C the values of R,, for the two surfaces are in good agreement, but at 500 "C the value for the KC1-coated vessel is noticeably higher. Nalbandyan16 has reported the detection of HO, radicals in the gas phase when H,O, is decomposed on a KC1-coated surface. Such a possibility would increase the concentration of HO, above the value calculated from the present mechanism. This 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 where the rate is high, reaction (7) becomes the main process removing HO, radicals and regeneration of HO, radicals by surface decomposition of H20, might become important.Unfortunately, no indication appears available from Nalbandyan's work of the yield of HO, radicals produced in this way. It is concluded that the KC1 results at the higher temperatures are too high, and that the Nalbandyan effect is the possible cause. Certainly the error is not due to the removal 4-298 Arrhenius Parameters for HO, + C,H, Table 3. Comparison of optimum values of R,, in boric acid- and KC1-coated vessels Rlo/(dm3 mo1-I s-l)' T/"C boric acid vessel KC1 vessel -~ 400 0.1 18 420 0.194 440 0.263 0.274 470 0.473 0.610 500 0.810 1.09 2.18 520 - - ~- of C,H, by OH radicals because this process is considerably more important in boric acid-coated vessels. Arrhenius Parameters for Reaction (10) Until relatively recently, a value of k , = 2.0 x lo9 dm3 mol-l s-l at 300 K, based on the results of several studies,, was accepted.This, together with Tree's" value of ca. 1.5 x lo9 dm3 mol-1 s-l at 1100 K from shock-tube studies suggested E, = 0 and a simple bimolecular H atom abstraction mechanism for reaction (7). Unfortunately, the mechanism is considerably more complex, at least at low temperatures, and the current position and views have been summarised recently by Kaufman and S h e r ~ e l l , ~ Patrick et al." and by Kircher and Sander.19 Only a brief outline will be given here. Several studies around 300 K have shown that the overall rate constant is pressure dependent,Ig 22 and increases markedly in the presence of polar gases such as NH, and H,0.4 Several studiesfg? 2 o t 23* 24 which together cover the range 240-510 K show clear evidence of a negative temperature coefficient for reaction (7).The most comprehensive determinations have been carried out by Sander and coworker^^^^ 22 who measured the HO, concentration by flash photolysis-u.v. spectro- scopy. For a number of diluent gases, the overall value of k , increased in a linear fashion with pressure, but with a marked intercept at zero pressure. For example, for N, at 298 K, k , increases from (1.15 k 0.17) x lo9 at 100 Torr to (1.79 0.27) x lo9 dm3 mol-1 s-l at 700 Torr. Unfortunately, data were not obtained below about 100 Torr, but extrapolation to zero pressure gives a value of k , = 1.05 x lo9 at 298 K.Within experimental error, the same zero pressure value was obtained when the diluent was Ar or SF,. Kircher and Sanderfg studied the effect of pressure (ca. 100-700 Torr) ever the range 240417 K and give Arrhenius parameters for the bimolecular and termolecular components of the overall rate constants. For Ar k,, = 1.38 x los exp [(600 f 130)/ TI dm:' mol--I s-' k,t = 3.05 x lo* exp [( 1 100 k 300)/ TI dm6 rnol-' s-' k,, = 1.32 x lo8 exp [(620 & 60)/ TI dm3 mol-1 s-l k,t = 6.9 x lo8 exp [(980&200)/T] dm6 rnol-, s-l. and for N, The following mechanism is suggestedfg to account for the results: b HO, + HO, 5 H20z 3 H,O, + 0, dM H,O, -+ products.R. R. Baldwin, C.E. Dean, M . R. Honeyman and R. W. Walker 99 With the steady state assumption, then the overall rate constant is given by k 7 = ka(kb + k c [ M 1 ) / ( k - ~ + k b + kc[M1)* The value of k , at zero pressure gives the bimolecular component k,, as and at low pressures (up to 700 Torr with N, and Ar)19. 22 where k,[M] -6 (kQ + k b ) then the termolecular component k,, is given by in agreement with the effect of pressure on the overall rate constant observed by Sander and 22 The present work was carried out at a total pressure of only 15 Torr, so that if this mechanism applies at high temperatures the values of k i , and k,, given above give k7 z k , , z 3.0 x lo8 dm" mol-1 s-l at 450 "C, the mean temperature used. However, this value is about a factor of five lower than Troe's17 value of 1.5 x lo9 at 1100 K.Use of Troe's value at 450°C would give a figure for k,, ca. t 5( =: 2.2) higher than if Sander's values are used for k,. If Troe'sI7 value at 1100 K is accepted, the most likely explanation is that it refers to the direct bimolecular step HO, + HO, + H,O, + 0, (7 d ) which would be expected to occur at a reasonable rate at high temperatures. To examine this explanation further, it is necessary to obtain an estimate of E,,. Since only Troe's single-temperature value is available at high temperatures, an estimate of A,d is required. Assuming simple Arrhenius behaviour, use of In ( A i d / L i d ) = AS,d/R with"> AS,, = 14.6 J K-l mol-1 at 500 "C gives A 7 d / A - , d =- 5.8. No experimental values of Aid or A - , , are available, but AWid is unlikely to be significantly different from A , , = 2.0 x lo1", determined from studies of the oxidation of HCHO over the temperature range 380-540 0C.26 With this value for then A , , = 3.4 x lo9, and combination with Troe's value of k , at 1100 K gives E,, = 6.7 kJ mol-l if, as indicated later, the low-temperature mechanism makes a negligible contribution to k , at 1100 K.HCHO + 0, + HCO + HO, (15) With these Arrhenius parameters, the direct bimolecular component k , , makes only about 1004 contribution to the overall value of k , at 300 K, so that the low temperature mechanism dominates, leading to a negative activation energy for reaction (7) at low temperatures. However, the Arrhenius plots for k,, obtained by Kircher and Sanderlg for both N, and Ar as diluents and for the equivalent rate constant for the mutual reaction of "0, radicals are noticeably curved with the higher temperature values of kib being higher than expected [see fig.5 for step (7b)l. A reasonable explanation of the curvature is that reaction ( 7 4 makes a significant contribution to the overall value of ki at the higher temperatures used by Sander and Kircher. If this is the case then the magnitude of 4 6 has been underestimated. Use of = 3.4 x lo9 dm3 mol-1 s-l and E,, = 6.7 kJ mol-1 permits the correction of the values of k,, obtained by Sander and Kircher for the contribution from step ( 7 4 . The results are shown in fig. 5. The correction at 240 K is only 4%, but rises to over 50"/, at 41 7 K, the temperature limit of the low temperature study.The corrected points now fit a good straight line with The nature of the assumptions and approximations made in the above analysis of reaction (7) and the uncertainties in some of the experimental data render the Arrhenius parameters suggested for steps (7b) and ( 7 4 liable to significant possible error. Nevertheless, the overall mechanism suggested and the parameters derived are reasonable and are recommended at this stage until further experimental evidence is available. Eib = - 10.1 kJ m01-l and A , , = 1.25 X 10, dm3 mo1-l S-'.100 Arrhenius Parameters for HO, + C,H, lo3 KIT Fig. 5. Variation of k, with temperature. x , results from Tree;" A, (N,); 0, (Ar) values of k,, given by Kircher and Sander;lg A, (N2), 0, (Ar) corrected values of k,, (see text).Dashed line: variation of k,, with temperature. Curve: calculated value of k , from k , = k,,+k,, (see text). At zero pressure, therefore, the overall value of k , is given by k , = k7b+k7d, and the curve drawn in fig. 5 shows the calculated value of k , as a function of temperature using the Arrhenius parameters derived above. The curve has a minimum at ca. 400 K, and if the suggested mechanism is correct it is clearly unfortunate that Sander and Kircher did not extend their study to at least 500 K. Pilling and Patrickz4 did study the reaction at higher temperatures (298-510 K), but at a fixed total pressure of N, of 700 Torr. Because of the importance of the termolecular component at this pressure, a minimum in the value of k, would be predicted at the slightly higher temperature of ca.460 IS. Pilling and Patrick suggest a simple negative temperature coefficient through the experimental values of k,, but the decrease on an Arrhenius plot is not uniform, and points between 380 and 455 K are virtually constant which could be due to a minimum in the value of k , in this region. Unfortunately, the single point at a higher temperature (510 K) is considerably lower, but it is also markedly lower than the mean line drawn by Patrick and Pilling and could possibly be seriously in error. For pressures above a few Torr, the termolecular component should be added into the overall expression for k,, but at combustion temperatures ( > 600 K) this component will only be significant at pressures well above atmospheric pressure.The values of k,,,/kjR. R. Baldwin, C . E. Dean, M . R. Honeyman and R . W. Walker 101 in the present work were obtained at a total pressure of 15 Torr, so that k , may be calculated solely from k , = k,,+k,,. At low pressures between 400 and 500 "C, E, = 5.8 kJ mo1-I and A, = dm3 mol-l s- l. Use of the recommended parameters for A l o / A $ and E,,-!J?, given earlier yields A,, = 1010.23*0.16 dm3 mol-1 s-l and E l , = 85.6k2.5 kJ mol-l. No other kinetic data are available for reaction (lo), but the parameters may be compared with A,, = 1010-38 dm3 mol-1 s-l and El6 = 84.5 8 kJ mol-1 for the reaction of HO, radicals with TMB, obtained under similar conditions and temperature range from direct measurements of k16/k!. Similar Arrhenius parameters would be expected as both alkanes contain only primary C-H bonds.A direct comparison is best made by comparing k,,/k$ with k,,/k$ at 450 "C, where values of 0.554 and 0.329 (dm3 mol-l s-l);, respectively, are obtained. On a per C-H bond basis, the value for TMB is a factor of 1.8 lower than that for ethane. An equivalent factor of 2.1 has been observed15 for OH attack on the two alkanes. As the ratios are effectively the same for the two radicals, despite the large difference in their selectivity, it is probable that steric effects influencing A factors are the main cause of the non-additivity observed HO, + TMB -+ H,O, + C8H17. Use of standard thermochemical data25. 26 gives ASlo = 21 J K-l mol-1 at 450 "C, so that with the assumption of simple Arrhenius expressions for (10) and (- 10) then A l 0 / A - , , = 12.5 and Apl0 = 1.4 x lo9 dm3 mo1-1 s-l.Similarly the value of AHlo = 50.5 kJ mol-l, so that El, = 35 kJ mol-l. These derived parameters for reaction (- 10) appear very reasonable when compared with those for H-abstraction reactions by ethyl radicals from compounds similar to H,0,.27 This work was supported by an S.E.R.C. research grant. References 1 L. J. Kirsch, paper presented at the Autumn Meeting of the Chemical Society, University of Hull, September 1984 (Symposium on Combustion Chemistry in the Gas Phase). 2 W. J. Pitz, C . K. Westbrook, W. M. Proscia and F. L. Dryer, Proc. 20th International Symposium on Combustion (Ann Arbor, August, 1984), to be published. 3 D. B. Smith, paper presented at the Autumn Meeting of the Chemical Society, University of Hull, September 1984 (Symposium on Combustion Chemistry in the Gas Phase).4 M. Kaufman and J. Sherwell, Prog. React. Kinet., 1983, 12, 2. 5 R. R. Baldwin, R. W. Walker and M. W. M. Hisham, J. Chem. SOC., Faraday Trans. 1, 1982,78, 1165. 6 R. W. Walker, Robert W. Walker and R. R. Baldwin, J . Chem. SOC., Faraday Trans. I , 1980, 76, 825. 7 R. W. Walker, in Reaction Kinetics (Specialist Periodical Report, The Chemical Society, London, 1975), 8 W. G. Alcock and B. Mile, Combust. Flame, 1975, 24, 125. 9 R. R. Baldwin and R. W. Walker, Proc. 17th Int. Symp. Combustion (The Combustion Institute, vol. 1, p. 161. Pittsburgh, 1979), p. 819. 10 R. R. Baldwin and J. A. Howarth, J. Chem. SOC., Faraday Trans. 1. 1982, 78, 451. 11 G. Atri, R. R. Baldwin, G. A. Evans and R. W. Walker, J . Chem. SOC., Faraday Trans. I , 1978,74,366. 12 R. R. Baldwin, I. A. Pickering and R. W. Walker, J . Chem. SOC., Faradaj. Trans. I , 1980, 76, 2374. 13 R. R. Baldwin, P. Doran and L. Mayor, Trans. Faraday SOC.. 1960, 56, 93. 14 R. R Baldwin, D Brattan, B. Tunnicliffe, R. W. Walker and S. J. Webster, Combust. FIume, 1970, 15 R. R. Baldwin and R. W. Walker, J . Chem. SOC., Faraday Trans. I , 1979, 74, 140. 16 A. B. Nalbandyan, Proc. 17th Int. Symp. Combustion (The Combustion Institute, Pittsburgh, 1979), 17 J. Troe, Ber. Bunsenges. Phys. Chem., 1969, 73, 946. 18 R. Patrick, J. R. Barker and D. M. Golden, J . Phys. Chem., 1984, 88, 128. 19 C. C . Kircher and S. P. Sander, J . Phys. Chem., 1984, 88, 2082. 20 R. A. Cox and J. P. Burrows, J . Phys. Chem., 1979, 83, 2560. 21 B. A. Thrush and J. P. T. Wilkinson, Chem. Phys. Lett., 1979, 66, 441. 22 S. P. Sander, M. Peterson, R. T. Watson and R. Patrick, J. Phys. Chem.. 1982, 86, 1236. 15, 133. p. 533.102 Arrhenius Parameters for HO, + C2H6 23 B. A. Thrush and G. S. Tyndall, Chem. Phys. Lett., 1982,92, 232. 24 R. Patrick and M. J. Pilling, Chem. Phys. Lett., 1982, 91, 343. 25 JANAF Thermochemical Tables, NSRDS-NBS 37 (U.S. Department of Commerce, National Bureau 26 R. R. Baldwin, G. R. Drewery and R. W. Walker, J. Chem. SOC., Faraday Trans. I , 1984, 80, 2827. 27 Handbook of Bimolecular and Termolecular Gas Reactions, ed. J. A. Kerr (CRC Press, Boca Raton, of Standards, Washington, 1971). Florida, 1981), vol. 1. Paper 5/417; Received 12th March. 1985

ISSN:0300-9599    

DOI:10.1039/F19868200089

出版商:RSC    

年代:1986

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1986, 82, 103-108 Homologation of Small Alkanes on Pt, Pd and Ni Catalysts Contribution of Intermediate Carbenes to Skeletal Isomerisation Antal Sarkany Institute of Isotopes of the Hungarian Academy of Sciences, P.O. Box 77, H- 1525 Budapest, Hungary Chain lengthening homologation of propane,n-butane,2-methylpropane,2- methylbutane and 2,2-dimethylpropane has been investigated over Pt, Pd and Ni catalysts. With both Pd and Ni the chain lengthening selectivity was observed to be commensurable to the isomerisation one. Addition of CH, onto the j?-carbon atom in n-butane and propane as well as chain lengthening of 2-methylpropane and 2-methylbutane were observed. The mechanism of the skeletal rearrangement is discussed in terms of CH, addition. _ _ _ _ _ ~ Under hydrogen-deficient conditions, the reaction of saturated hydrocarbons over transition metals has been observed to be accompanied by the formation of hydrocarbons with higher carbon number than the reactant.l The chain lengthening reaction (homo- logation) detected on a large group of metals4$ has been interpreted by the recombination of intermediate carbenes with terminal olefins uia the formation of a metallocyclobutane ring.3* At elevated temperatures the reaction of C, and C, alkanes with Ni59 ‘ 9 * and Pd4 catalysts, if the chain lengthening is not hindered by branching, results in the formation of benzene and alkylaromatics.This indicates that the addition of carbene to the non-substituted end of intermediate alk- 1 -ene is immediately followed by 1,6- dehydrocyclisation on surfaces poisoned strongly by carbonaceous deposit^.^^ In contrast to these findings, the preliminary results in our laboratory’ and those presented by O’Donohoe et aL4 have shown that the chain lengthening with small alkanes, although the selectivity is not too high and drops with temperature, yields iso-alkane and n-alkane in commensurable quantity. Moreover, the formation of methylcyclopentane from cyclopentane has been observed over some Cu-Ni/SiO,, Ag-Ni/SiO, and Au-Ni/SiO, catalysts,’ suggesting that the internal double bonds may become accessible to carbenes under certain experimental conditions.In principle, a CH, addition-abstraction mechanism, in which the CH, unit is added to a /?-carbon atom in n-alkane, might lead to the formation of the isomer of the parent hydr~carbon.~ As few data are available on the mechanism of isomerisation at low excess of H, we decided to undertake a study on the selectivity of carbene insertion onto butane, propane and 2-methylpropane over Ni, Pd and Pt catalysts.Experimental The experiments were accomplished in a glass circulation system of 0.1443 dm3 attached to a gas chromatograph. The experimental set up and the analysis of the products have already been de~cribed.~ The transformation of hydrocarbons was measured up to 20-30% conversion and the rate of the formation of the products was used for the calculation of product selectivity. The selectivity is defined in accordance with earlier publicati~ns.~q 9 9 lo The catalysts have been prepared by a conventional technique.1 wt % Pd/SiO, and 5 wt % Ni/SiO, catalysts were prepared by the method of incipient 103104 Homologation of Alkanes on Pt, Pd and Ni wetness using Pd(NH,),Cl,H,O and Ni(NO,), solutions, respectively. The catalysts were dried and calcined at 673 K for 5 h and reduced in a stream of H, at 693 K for 20 h. Prior to the catalytic and chemisorption measurements, the samples were stabilised by rt:peated oxygen-hydrogen treatments at 623 K. The dispersion, along with the experimental conditions, are presented in the tables. The preparation and characteristics of the Pt catalysts have been presented elsewhere.ll Results and Discussion The results with Pt/SiO, and Pt-black, 5 wt % Ni/SiO, and 1 wt % Pd/Al,O, catalysts are summarised in tables 1-3, respectively. Since the objective of the present study was to investigate chain lengthening, most of the hydrocarbon reactions were carried out at low H,/HC ratios. With Pt catalysts, even under these conditions, a considerable amount of 2-methylpropane was formed from n-butane and, in agreement with the previous findings,*v5 the selectivity for the formation of the next higher alkane was low.Nevertheless, both 2-methylbutane and n-pentane could be detected without any difficulty. The isomerisation selectivity of the highly dispersed 1 wt % Pt/SiO, sample is superior to that of Pt-black and the rupture of the C-C bond in n-butane is almost uniform with the former catalyst. Both the high isomerization selectivity and the low excess of methane formation over 1 wt % Pt/SiO, can be interpreted by the prevalence of localised Pt-hydrocarbon interactions.ll Over 5 wt % Ni/SiO, and 1 wt % Pd/Al,O, catalysts the selectivity of isomer forma- tion and that of chain lengthening were observed to be commensurable.First, we consider the selectivity data with respect to chain lengthening. The inspection of the results with n-butane in tables 2 and 3 shows that the selectivity for 2-methylbutane is roughly twice that for n-pentane, suggesting that the addition of a C, unit to a P-carbon atom is favoured under these conditions. This result is also confirmed with propane over 5 wt % Ni/SiO, catalyst. One could argue that 2-methylbutane might be formed by the immediate isomerisation of an intermediate n-pentane. This suggestion could only be accepted if the isomerisation selectivity of Ni and Pd were high enough, which is not the case as shown by the penultimate column in tables 2 and 3.The experiments with 2-methylpropane also allow us to reject this suggestion as the ratio of n-pentane over 2-methylbutane is very small (0.05 1 and 0.03 1 on Pd/Al,O, and 0.07 1 on Ni/SiO,). While the formation of 2-methylbutane could be defined by the addition of a CH, group to 2-methylpropane, n-pentane might be formed either by bond-shift skeletal isomerisation of 2-methylbutane or, more likely, by consecutive CH, addition-abstraction steps: The results with 2-methylpropane and 2-methylbutane, on the other hand, provide evidence that the presence of a methyl substituent significantly inhibits the probability of chain lengthening.On our highly dispersed Ni/SiO, sample (63% dispersion) the W h o m / Whydr ratio (the rate of homologation compared with the rate of hydrogenolysis) is observed to be 11.78 x lo-, and 3.01 x with n-butane and 2-methylpropane, re- spectively, under comparable experimental conditions. As shown by the data in table 3, the difference between the selectivity values measured with n-butane and 2- methylpropane is even less on the 1 wt % Pd/A1,0, catalyst. A further check was made with 2-methylbutane over Ni/SiO, (table 2, last two rows): the selectivity of the formation of 2-methylpentane is close to that of 3-methylpentane. This latter compound is likelyTable 1. Transformation of n-butane over 1 wt % Pt/Si02a and Pt-blackb selectivity (% ) catalyst 1 wt % Pt/SiO, 557 1.33 6.65 9.89 16.10 8.91 82.21 0.108 0.063 4.76 0.128 1 wt % Pt/SiO, 555 3.99 4.52 13.95 18.41 13.13 75.79 0.244 0.306 3.36 3.05 Pt-black 556 1.33 6.25 43.28 11.91 25.61 63.81 0.157 0.473 1.77 2.19 Pt- blac k 556 3.99 4.39 42.14 20.80 35.26 50.67 0.713 0.832 1.07 4.07 a Dispersion = 80%.B.E.T. surface area = 5.4 m2 g-l. The rate of homologue formation (iC5+nC,) is transformed to C, equivalents. b Table 2. Transformation of alkanes on a 5 wt % Ni/SiO, catalysta* 2 < selectivity c') it .~ 1ooFqso lOOKom K ydr K y d r HC T/K p(H,)/kPa C, C, C, iC, nC, iC, nC, 2M2Bu2- 2MP 3MP nH MCP B ~ Pr nBu nBu 2MPr 2,2DMPr 2,2DMPr 2MBu 2MBu 556 4.39 193.0 51.5 551 4.52 161.1 21.6 552 4.39' 291.9 21.3 556 4.39 249.3 19.7 553 4.65 454.8 12.1 513 4.65' 154.1 14.0 554 6.18' 155.6 11.1 553 4.65 113.1 9.3 ~ ____ - - - - __ 2.17 1.83 - - - - - 42.5 7.2 - 4.25 2.60 - ~ ~ 0.86 24.5 2.9 - 0.7 0.31 - - - 26.3 - 4.5 2.1 0.14 - __ - - - - - - - - - - - - - - -~~ 2.1 3.6 - 31.2 56.0 - 12.2 28.7 33.0 - 1.41 4.1 0.36 0.36 - - - - - - - - - - - 1.33 8.7 24.1 29.8 - 0.72 19.8 1.52 0.81 0.17 0.98 0.93 - 5.4 8.65 11.78 2.91 1.22 4.81 3.01 0.00 0.00 0.00 0.00 I .24 2.68 0.96 7.12 ~~ a Dispersion = 63%.Pr, nBu, 2BMPr, 2,2DMPr, 2MBu, 2M2Bu2-, 2MP, 3MP, nH, MCP and B are propane, n-butane, 2-methylpropane, 2,2-dimethylpropane, 2-methylbutane7 2-methylbut-2-ene7 2-methylpentane, 3-methylpentane, n-hexane, methylcyclopentane and benzene, respectively. Measurements with 1.33 k Pa hydrocarbon, otherwise with 3.99 k Pa.% 0 Table 3.Transformation of n-butane, 2-methylpropane and 2,2-dimethylpropane on 1 wt % Pd/Al,03 catalysta ~- selectivity (% ) nBu 524 2.66 9.84 93.6 2.6 87.8 3.21 - 3.99 4.79 64.5 3.8 78.0 8.74 - nBu 553 nBu 558 1.33 6.31 87.1 5.9 84.8 8.92 - 2MPr 523 2.66 9.57 100.2 1.5 90.4 - 5.44 2MPr 553 3.99 4.52 65.9 1.9 72.4 - 15.43 2,2DMPr 553 3.99 4.39 111.3 3.6 28.1 74.2 tr ~ - ~ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ ~ _ ~ _ _ _ a Dispersion = 8.5%. 3.63 1.41 3.53 5.56 7.67 4.06 1.42 19.15 1.35 0.78 10.10 2.42 2.51 0.13 5.8 1 3.52 9.84 0.3 1 21.48 14.13 0.00 0.00 0.00 0.00 ~~ ~ - ~~A . Sarkany 107 to be formed by the recombination of a CH, unit with the branched end of 2-methyl- butane. All the observations in this study seem to indicate that in small alkanes the insertion of CH, to a /?-carbon atom is not hindered and at the same time the small isoalkanes are also effective in the chain lengthening.With C, and C7 n-alkanes the addition of CH, to the internal carbon atoms was observed only to a very limited extent over Ni and Pd catalysts.4* 5 9 The disagreement between the results of this paper and those reported in ref. (4) and (7) might be resolved if one accepts that with small alkanes, by virtue of their less effective carbonpoisoning activity, the surface sites are sterically crowded to a smaller extent than those with large alkanes. The small size of the incoming hydrocarbon and a less effective steric crowding might promote both the formation and the chain lengthening of an iso-alkane. One cannot rule out, however, that in the previous high temperature experiment^,^?^ with C,-C, hydrocarbons over Ni and Pd, that the rapid aromatisation and the relatively high chemical stability of benzene have obscured the real selectivity of carbene addition. The observed results afford some insight into the mechanism of isomerisation over Ni and Pd catalysts.There has been much discussion in the literature concerning the mechanism of the bond-shift type skeletal rearrangement. The experiments with caged hydrocarbons12 in excess hydrogen and deuterium labelling l7-l9 have provided strong evidence that a monoadsorbed intermediate (a-alkyl adsorbed radical) is sufficient for bond-shift skeletal rearrangement (Rooney-Samman mechanism). On the basis of the experimental results in this paper, we now have good reason to suggest that, under ‘ hydrogen deficient ’ conditions, intermediatecarbenes also contribute to the isomerisation process.Apparently, under these reaction conditions, the H,/HC ratio will decide which reaction route becomes conspicuous. In principle a large excess of hydrogen should favour the Rooney-Samman mechanism, whereas the high reaction temperatme and low excess of hydrogen might promote a reaction route viu participation of intermediate carbene. A carbene route for isomerisation has been suggested, but not proved, by Gault and Muller16 in their dehydrocyclisation studies on Pd : The elementary been discussed rearrangement, steps of the ‘ carbene metathesis ’ route (Garin-Gault mechanism) have in detail in ref. (20)-(22).The key step in this intramolecular-type as emphasised by Gault and Garin, is the rotation of an intermediate alkene in the vicinity of a reactive CH, group. However, the rotation or the migration of a surface olefin might lead to the separation of the CH, unit from the surface alkene, so that the intramolecular process might become an intermolecular one. The rotation step of the ‘carbene metathesis’ route has already been questioned by Ponec.15 Considering the results in tables 2 and 3 we suggest that the formation of isomers under these experimental conditions is part of the homologation process. A CH, addition- abstraction route (the elementary processes are not depicted) is presented :108 Homologation of Alkanes on Pt, Pd and Ni An alternative route, i.e.a CH, abstraction-addition mechanism, cannot be ruled out entirely as the chain lengthening with propane yielded both n-butane and 2-methylpropane. It should be noted, however, that in the hydrogenolysis of the C, hydrocarbons, say on Ni, the surface concentration of the C , intermediate alkene might not be high enough, owing to the propensity for deep fragmentation, whereas the interaction of the reactant C, alkane with poisoned sites would produce a sufficient amount of intermediate alk-1-lene. The last step of the above scheme, i.e. the fission of a CH, group, requires some discussion. As shown by results in table 2, starting from 2-methylpentane there is an almost equal chance of forming n-butane and 2- methylpropane. The intermediate 2-methylpentane, once formed, might either desorb and chemisorb from the gas phase (the reactions were followed up to 2630% conversions), or it might remain on the surface and lose one of its terminal carbon atoms.Considering the results with 14C-labelled hydrocarbons over Ni catalysts,s- 23j 24 one might propose that the C , alkanes are formed without the desorption of intermediate 2-methylpentane. The absence of the CH, insertion on to 2,2-dimethylpropane explains why the isomerisation of this hydrocarbon was not observed over Ni and Pd. Under hydrogen- deficient conditions the CH, addition-abstraction method of isomerisation might contribute to a small extent to the formation of isomers on Pt catalysts, but, as has already been empha~ised,~~ l4 the homologation reaction route is not a preferred one on this metal.References 1 J. R. Anderson and B. G. Baker, Proc. R. SOC. London, Ser. A, 1963, 271, 402. 2 A. Peter and J. K. A. Clarke, J . Chem. SOC., Faraday Trans. I , 1976, 72, 1201. 3 C. O’Donohoe, J. K. A. Clarke and J. J. Rooney, J . Chem. Soc., Chem. Commun., 1979, 648. 4 C. O’Donohoe, J. K. A. Clarke and J. J. Rooney, J. Chem. Soc., Faraday Trans. I , 1980, 76, 345. 5 A. Sarkany and P. TetCnyi, J. Chem. Soc., Chem. Commun., 1980, 525. 6 A. Sarkany, S. Palfy and P. Tttenyi, React. Kinet. Catal. Lett. 1980, 14, 345. 7 A. Sarkany, S. Palfy and P. TetCnyi, Acta Chim. Hung. Acad. Sci., 1982, 111, 633. 8 A. Sarkany, J. Catal., 1984, 89, 14. 9 L. Guczi, A. Sarkany and P. Tetenyi, J . Chem. Soc., Faraday Trans. 1, 1974, 70, 1941. 10 V. Ponec and W.M. H. Sachtler, in Proc. 5th Int. Congr. Catal., ed. J. Hightower (Miami Beach, 1972), 1 1 A. Sarkany, J. Gaal and L. Toth, in Proc. 7th Int. Congr. Catal., ed. T. Seiyama and K. Tanabe 12 J. K. A. Clarke and J. J. Rooney, Adv. Catal., 1976, 25, 125. 13 0. Zahraa, F. Garin and G. Maire, Faraday Discuss. Chem. Soc., 1981, 72, 45. 14 J. J. Rooney, Faraday Discuss. Chem. SOC., 1981, 72, 87. 15 V. Ponec, Faraday Discuss. Chem. Soc., 1981, 72, 88. 16 J. M. Muller and F. G. Gault, J. Catal., 1972, 24, 361. 17 Z. Karpinski and L. Guczi, J . Chem. SOC., Chem. Commun., 1977, 563. 18 Z. Karpinski, Nouv. J. C; ‘yz., 1980, 4, 561. 19 0. E. Finlayson, J. K. A. darke and J. J. Rooney, J. Chem. Soc., Faraday Trans. I , 1984, 80, 191. 20 F. G. Garin and F. G. Gault, in Chemistry and Chemical Engineering of Catalytic Processes, Ser. E, Applied Sciences, no. 39, ed. R. Pnns and G. C. A. Schuit (Sijthoff and Noordhoff, Alphen aan den Rijn, The Netherlands, 1980), p. 351. 21 F. G. Gault, Adv. Catal., 1981, 30, 1 . 22 G. L. C. Maire and F. G. Garin, in Catalysis Science and Technology (Springer, Berlin, 1984), vol. 6, p. 161. 23 L. Guczi, A. Sarkany and P. TCtenyi, in Proc. 5th Int. Congr. Catal., ed. J. Hightower (Miami Beach, 1972), vol. 1, p. 1 1 1 1 . 24 A. Sarkany, L. Guczi and P. Titenyi, Acta Chim. Hung. Acad. Sci., 1975, 84 245. vol. 1, p. 645. (Publisher, Tokyo, 1980), vol. 1 , p. 291. Paper 51432; Receked 15th March, 1985

ISSN:0300-9599    

DOI:10.1039/F19868200103

出版商:RSC    

年代:1986

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I, 1986, 82, 109-118 Dynamic Behaviour of Fluorescence Quenchers in Cetyltrimethylammonium Chloride Micelles Angelos Malliaris,? Jacques Lang and Raoul Zana* C. R. M . , Grkco Microkmulsions, CNRS, 6 rue Boussingault, 67000 Strasbourg, France The kinetics of quenching of the fluorescence of pyrene solubilized in cetyltrimethylammonium chloride (CTAC) micelles by three different types of quenchers [pyrene or cetylpyridinium chloride (immobile quenchers), dodecylpyridinium chloride (mobile, hydrophobic quencher) and iodide ion (mobile, hydrophilic quencher)] has been investigated by means of time- resolved fluorescence measurements. The results obtained with the two immobile quenchers have permitted the first determination of the mean micellar aggregation number N of CTAC at high concentration up to 1.2 mol dmP3.These N-values have then been used for a quantitative evaluation of the kinetic and equilibrium quenching parameters for the other two types of quenchers. The results concerning the quenching by the iodide ion are discussed on the basis of the available models proposed for this type of quenching. They appear to favour the electrostatic model of Almgren et al. The kinetics of quenching of the fluorescence of micelle-solubilized probes by various quenchers both organicl-s and inorganic7* 9-13 has been recently the subject of extensive investigations. Taking into account the fact that the hydrophobic fluorescent probes such as pyrene or methylpyrene used in most of these investigations have a residence time inside the micelle much longer than their fluorescence lifetime, the mechanism of the quenching process depends exclusively on the nature of the quencher. Thus three situations can be distinguished.(i) The distribution of the quencher is ‘ frozen on the time scale of the fluorescence, i.e. when the residence time of the quencher in the micelles is much longer than the probe fluorescence lifetime. This is the case of pyrene excimer for ma ti or^^^^^^ where the pyrene fluorescence is quenched by pyrene in the ground state, or the quenching of pyrene fluorescence by cetylpyridinium i o n P (CP;), the so-called ‘immobile’ quenchers.l79 l8 (ii) The quencher is hydrophobic and partitioned between the micelle interior and the aqueous phase. The quencher may then be exchanged between two micelles in a time comparable to the fluorescence lifetime. This is the case of the ‘ mobile ’ quenched7 7 l8 such as m-dicyanobenzene,8 and dodecylpyridinium ion (DPC).(iii) The quencher is hydrophilic and ionic such as I-, Cu2+, etc. and is partitioned between the micelle Stern layer and the aqueous phase. This case is by far the most complex7~ l3 and it is observed that the rate of intermicellar quencher exchange increases rapidly with the surfactant concentration.13 Two different interpretations have been given to this observation. In the first one this increase is attributed to ‘close encounters’ of micelles during which the quencher effectively ‘hops’ from micelle to mi~elle.~O-~~9 l9 The rate constants for the association of the quencher to a micelle, and its dissociation from a micelle, k+ and k-, respectively, are assumed to remain constant and the increase in the exchange rate with increasing surfactant concentration C (and thus micelle concen- tration [MI) then results from the increase of the number of close encounters.The second interpretation is based on a numerical solution of the Poisson-Boltzmann equation using a cell model for the micellar so1ution.203 21 These calculations show that a large increase t On leave of absence from N.R.C. ‘Demokritos’, Athens, Greece. 109110 Fluorescence Quenching in CTAC Micelles in k- with C results from the decrease of the electrostatic potential at the micelle surface as [MI increases, which facilitates the release of the counterions located in the micelle Stern layer.In the present study we have investigated the quenching of pyrene in aqueous micellar solutions of cetyltrimethylammonium chloride (CTAC) up to very high surfactant concentration (1.2 mol dm-3, i.e. just below the range of liquid crystal phases) by means of two immobile quenchers, pyrene and CP;, a mobile hydrophobic quencher DP;, and the hydrophilic ionic quencher I-. The objective of this work was to investigate the quenching mechanisms for the various types of quenchers, with emphasis on ionic quenchers, particularly at the high [MI range where intermicellar quencher exchange becomes very important.20* 21 Moreover we have examined the possibility of obtaining the mean micelle aggregation number N and rate constants of the quencher-micelle association-dissociation process, for the various types of quenchers.Finally we have used our experimental data to test the various assumptions most often adopted for the extraction of micellar parameters from quenching experiments. 7 , l3 CTAC was selected for this study, because (i) it is easily prepared and purified, (ii) it has a Krafft temperature which allows the use of concentrated solutions at 298.15 K and (iii) its micelles are dynamically stable on the time scale of fluorescence.22 Materials and Methods The sample of CTAC was obtained by exchanging the bromide ions of hexadecyltri- methylammonium bromide for chloride ions with an ion exchange resin (Merck 111). CTAC was purified by three crystallizations in ethanolkthylacetate mixtures.The c.m.c. of CTAC (1.38 x 1 0-3 mol dmP3 at 298.15 K) was determined by electrical conductivity using a Wayne-Kerr conductometer operating at 1.592 KHz. Dodecylpyridinium chloride (DP,C) was purchased from K and K and purified by boiling its methanol solution with carbon black and by several crystallizations in ace tone-e thylace ta te mixtures. Hexadecylpyridinium chloride (CP,C) was obtained by reacting pyridine with chloro-1-hexadecane in dry ethanol for 4 h at 408.15 K in an autoclave and purified by two crystallizations in ethylacetate and two crystallizations in acetone. Pyrene (Aldrich 99 % ) was purified by extensive zone refining. Potassium iodide (Prolabo > 99.5%) was used as received. Pyrene was dissolved into the micellar solutions following aclassical procedure.23 For the pyrene excimer experiments the [pyrene]/[micelle] molar concentration ratio was kept close to unity.For the measurements based on the quenching of pyrene fluorescence by DP;, CP; or I-, the [pyrene]/[micelle] and the [quencher]/[micelle] ratios were about 0.05 and 1, respectively. These conditions ensured negligible perturbation of the micelle by either the probe or quencher and accurate determinations of N . All micellar solutions were purged with pure argon in quartz cuvettes for 30 min. All measurements were performed at 298.15 K. The fluorescence decay curves of micelle-solubilized pyrene were obtained by the single photon counting technique24 and analysed using a non-linear weighted least-squares procedure. Theoretical The equations which govern the decay of a fluorescence probe in the presence of a quencher in a micelle have been given in several 10,12, 1 7 9 18, 2 5 9 26 We shall briefly recall the main equations and assumptions involved in order to facilitate the discussion of the results.In all situations, the fluorescence decay curves have been shown to obey the equation 2 5 3 2 6 I ( t ) = I(0) exp(-A,t--A,[l -exp(-A4t)]) (1)A . Malliaris, J . Lang and R. Zana 111 where I ( t ) and I(0) are the fluorescence intensities at time t and t = 0, following an excitation with a &light pulse. A,, A , and A , are constants whose expressions in the [QI k, k+ + KkJM] general case are given by:l09 A , = k,+- A , l+K[M] A4 = k, + kJM] + k- (4) where [Q] is the total quencher concentration, k, is the rate constant for the decay of the fluorescence of the micelle-solubilized monomeric probe, k , is the pseudo first-order rate constant for intramicellar quenching, k+ and k- are the rate constants for the association and the dissociation of the quencher to/from a micelle, respectively, K = k + / k - is the equilibrium association constant of the quencher to the micelle, and k, is the second-order rate constant for the exchange of quencher between micelles through micellar collisions or close encounters. Here it should be pointed out that through rearrangement of eqn (2)-(4) one obtains A3 k Q = A2-k0+A3A,' This equation allows the estimation of k , from the values of the experimentally available parameters.Therefore from eqn (4) and (5) k , and k,[M] + k- can be obtained directly from the fluorescence decay parameters for all types of quenchers, independently of any assumptions or approximations made.It has been shown17 that the time-resolved fluorescence method can be used to obtain information on micelle aggregation numbers only when k, 3 k,. The three situations discussed in the introduction can now be considered more quantitatively. (a) Immobile Quencher This corresponds to k,, k , % k- +k,[M]. Provided that K[M] % 1 , which is fulfilled for the quenchers used in the present work, eqn (2)-(4) then reduce to A , = k, A, = [QI/[Ml A , = k,. Analysis of the decay curve thus permits the calculation of the value of [MI and, in turn, of the mean micelle aggregation number N from N = (C-c.m.c.)/[M]. (9) (6) Mobile Hydrophobic Quencher If it is assumed that the intermicellar repulsions between ionic micelles nearly prevent collisions, then k,[M] 6 k- and eqn (2) and (4) reduce to k, k+ A, 1 + K[MJ A , = k,+- A , = k, + k- (1 1) whereas eqn (3) is unchanged.The main difference between the cases of mobile and immobile quenchers is that whereas A , is in both situations independent of [Q], A ,112 Fluorescence Quenching in CTAC Micelks becomes a linearly increasing function of [Q] at constant [MI in the case of a mobile quencher. This provides an easy test of the type of system under investigation. In any case it is important to realize that there are four unknowns, k,, k+, k- and [MI, for only three experimentally available quantities, A,, A, and A,. However, as was mentioned earlier, k , and k- can always be calculated from eqn (4) and (9, but in general it is not possible to obtain k+ and [MI, and thus N , from the data pertaining to the mobile hydrophobic quencher.Some additional assumption must therefore be made. Thus [MI can be assumed to be proportional to C-c.m.c., which is equivalent to assuming that N is independent of C. [MI is then obtained from the slope of the plots of A , uersus [Q] at different C. This method permits one to obtain both k+ and [MI. It can also be assumed that K[M] + 1, i.e. k+[M] + k-. On this assumption only [MI can be obtained from the results. It will be seen below that either assumption may not be valid in many instances. If no assumptions are made, k+ can be obtained from eqn (10) only if N has been determined in an independent experiment with an immobile quencher. (c) Mobile Hydrophilic Quencher The experimental results show that both A , and A , are increasing functions of [Q].If the intermicellar exchange of quencher is attributed only to the hopping of a quencher from micelle to micelle upon close encounter of micelles, then k,[M] % k-. This situation (pure hopping) has never been encountered in practice, but cannot be excluded. With hydrophilic ionic quenchers such as I- or Cu2+ it has been assumed that together with the hopping mechanism, the quencher is also exchanged ilia the aqueous phase and the full eqn (2) and (4) must be used to describe the system.7, R~ l3 k was considered to be a true rate constant, i.e. independent of both [Q] and [MI.In this case k , is still given by eqn ( 5 ) , and Thus, in principle both k , and k can be obtained from the plot of A , - k , L’ersus [MI provided [MI is known. This takes us back to the discussion at the end of the paragraph concerning the mobile hydrophobic quencher. If close encounters of ionic micelles are assumed to be prevented by intermicellar repulsions, the intermicellar exchange of quencher can only occur cia the aqueous phase. In the ‘electrostatic’ model developed by Almgren et al.z0.21 k- is shown to be an increasing function of [MI. Eqn (3), (10) and (1 1) describe this situation and k , and k- can be obtained at each surfactant concentration from eqn (5) and (1 l), provided [MI is known. There is however a basic difference with the case of the hydrophobic mobile quencher: k- as well as k+ are now dependent on [MI, and K is not a true equilibrium constant.Recall that the binding of counterions to micelles has the characteristics of a condensation 28 as in the case of polyelectrolyte solutions.29 This fact alone would exclude the use of an equilibrium constant to describe counterion binding to micelles. The situation is different for non-ionic solutes (such as fluorescence probes) or hydrophobic mobile quenchers such as DP:, the binding of which involves forces other than electrostatic forces. Both the ‘hopping’ model and the ‘electrostatic’ model predict qualitatively similar dependence of the quencher exchange rate on the micellar concentration at low [MI. However, at high [MI the electrostatic model requires a steep non-linear increase of exchange rate with [MI while the hopping mechanism always predicts a linear relationship between [MI and the exchange rate.The experimental evidence published thus fars* 21 deals exclusively with low [MI where both mechanisms predict similar behaviour and do not permit a choice between these mechanisms. Moreover since the important parameter is the micellar concentration [MI and not the overall surfactant concentration C it is imperative to know the change of N with C. In previous studies13 N was assumed to be independent of C, which introduced some error in the calculation of the rate constants of interest. This is avoided in the present work. k,[M]+k- = A , - k , . (12)A . Malliaris, J . Lang and R. Zana 113 Results and Discussion (a) Immobile Quencher (Pyrene Excimer Formation and Quenching of Pyrene by CP?) The values of A,, A , and A , are listed in table 1.It is seen that A , is independent of both C and [Q], and equal, within the experimental error, to k,, the rate constant of fluorescence decay of pyrene solubilized in CTAC micelles (k, = 2.94 x lo6 s-l was obtained in a preliminary experiment at a pyrene concentration of ca. lop5 mol drn-,, at which concentration excimer formation is negligible). These results indicate that pyrene and CP; can be considered as immobile quenchers. Eqn (7) and (8) have been used to obtain the micelle aggregation numbers and the quenching rate constants which are listed in table 1 under headings N and k,, and plotted in fig.1. The N-values obtained with pyrene only (excimer formation) and with the pyrene/CP$ system (quenching) are seen to fall on the same curve. The aggregation number shows a sizeable increase with C, with a nearly linear increase in a fairly large range of concentration between 0.2 and 1 mol dm-3. Above 1 mol dmP3, N increases rapidly possibly owing to the approach of the phase boundary limit. Aggregation numbers for CTAC have been reported for low surfactant concentrations. Values of 80-1 15 were found by fluorescence quenching measurements for CTAC concentrations 0.0065-0.05 mol dm-3.7r l4 These values agree, within experimental accuracy, with the ones reported here. On the other hand, a small aggregation number of 59 was found by e.s.r.,O for [CTAC] z 0.01 mol dm+.However, a discrepancy in the value of the c.m.c. compared to the one reported in the literature was also found by the same e.s.r. method (4 x c.f. 1.38 x lop3 mol dm-3). In any case this aggregation number is very small for a surfactant with a 16 carbon atom aliphatic chain31 and appears to be in error. To the authors’ knowledge there are no reported N-values at high C with which the present measurements can be compared. Although the value N = 154 obtained at C = 1 mol dm-3 is significantly larger than the value 94 which corresponds to the minimum spherical micelle for a surfactant with a 16 carbon atom alkyl chain,31 the shape of the CTAC micelle does not differ much from that of a sphere. Thus, if for the sake of discussion we assume a spherocylindrical shape, the ratio between the overall length and diameter of the micelle is found to be only ca.1.7 on the assumption of a surface area of 0.7nm2 per surfactant head group in the cylindrical part of the micelle. The near-spherical micellar shape at high concentration is further suggested by the k , data listed in table 1. Even though k , decreases significantly upon increasing C , and thus N , the quantity NkE is constant, within the experimental error, as predicted by theoretical treatments of intramicellar quenching kinetics in spherical micelle~.~~ Thus, our results show that, contrary to CTAB micelles, which become very elongated at C > 0.2 mol dm-3,33 CTAC micelles remain nearly spherical, even at concentrations as high as 1.2 mol drn-,. This conclusion provides additional evidence of the strong influence of the nature of the counterion on the micelle size and shape in the case of cationic surfactant s o l ~ t i o n s .~ ~ ~ 35 (b) Mobile Hydrophobic Quencher (DPT) The values of A,, A , and A , for the pyrene/DP,+ system are also listed in table 1. It is seen that the two systems having about the same surfactant concentration (0.47 mol drn-,), but largely differing [Q] values are characterized by different A , values and nearly equal A , values, as is expected from eqn (10) and (1 1) which apply to the case of a mobile hydrophobic quencher. The [MI values obtained in the case of the immobile quenchers have been used for the analysis of the data by means of eqn (3), (10) and (1 1) to yield the values of k,, k- and k+ listed in table 1 .It can be seen thatTable 1. Experimental parameters for quenching of the fluorescence of CTAC-solubilized pyrene -~ . ~ C [QI 4 A3 A4 kQ 4 - k , k+ K / lo5 s-I / lo9 dm3 mol-I s-l / lo3 dm3 mo1-l K[M] /mol dm-3 /I0 mol dm /I06 s-1 /I06 SS' / 106 s-1 Nb 0.016 0.03 1 0.064 0.123 0.254 0.300 0.465 0.740 0.991 1.204 0.12 0.205 0.46 0.73 0.051 0.193 0.458 0.499 0.017 0.050 0.052 0.052 0.052 0.100 0.246 0.303 0.505 0.840 1.36 2.64 5.17 8.15 16.7 30.5 25.9 38.3 51.0 44.9 7.9 12.7 26.0 37.1 3.74 12.2 64.9 29.1 1.74 3.67 2.79 3.96 5.15 6.7 15.97 17.16 28.6 49 3.04 3.03 2.94 2.90 2.93 3.06 2.99 3.01 3.18 3.01 3.01 3.03 3.05 3.01 3.1 3.2 3.4 3.16 3.45 3.52 3.44 3.65 3.8 3.73 4.44 4.52 5.29 6.12 0.84 0.88 0.87 0.79 0.82 1.30 0.75 0.74 0.79 0.7 0.80 0.74 0.77 0.76 0.51 0.66 1.62 0.72 0.95 0.71 0.54 0.7 0.87 0.63 0.49 0.42 0.3 0.17 pyrene excimer formation 8.4 8.4 7.9 7.9 6.8 6.8 6.1 6.1 5.7 5.7 5.7 5.7 5.2 5.2 5.0 5.0 4.4 4.4 3.6 3.6 quenching by CP; 10.2 10.2 10.0 10.0 8.8 8.8 8.4 8.4 quenching by DP; 12 11.8 11.5 11.2 9.15 8.9 8.7 8.5 8.77 7.94 8.7 8.9 8.9 7.8 9.3 8.9 9.5 13.4 quenching by I 8.3 7.3 8.0 8.1 8.0 6.8 7.I 6.3 5.3 5.5 91 99 106 118 124 128 134 143 154 187 I17 119 135 145 103c 70d 122 100 133 120 135 130 9Ic 89d 104 100 105 102 105 108 105 1!2 114 117 125 126 126 136 135 140 145 155 - ~ ~ - - ~ - _ _ ~ ~ ~ - 2.2 3.2 2.5 2.4 4.7 6 4 7.0 8.0 9.0 10.0 12.0 26.0 42.0 79.0 a Accuracy+4%. Accuracy$lO%. r N-values from the immobile quencher measurements. N-values from the corresponding mobile quencher on the assumption k[M] 9 1A .Malliaris, J. Lang and R. Zana 115 I I I I I I I I 1 ~ - 1 1 lo ~~ 0 0.2 0.4 0.6 0.8 1.0 1.2 C/mol dm-3 Fig. 1. Micellar aggregation numbers N for CTAC obtained by various quenchers (see text). Quenching by @, excimer; A, CP;; A, DP; and +, I-. 0, k , for excimer formation. k- = (2.7 kO.5) x lo5 s-l and k+ = (1.2 0.5) x lo9 dm3 mol-1 s-l. These values, how- ever, cannot be compared directly to those obtained from chemical relaxation interpreted on the basis of Aniansson and Wall theory3' for the exchange process, and which are k ~ , z 4.5 x lo7 s-l and k& 2 3 x lo9 dm3 mol-1 s-l. Indeed in chemical relaxation, k,, corresponds to the exchange of DP; between DP, C micelles and the bulk. Thus all the surfactant ions making up the micelles are exchanged whereas, in fluorescence, only probe exchange is observed and k- must be compared to hz-- / N e R Y ln5 c-1 Ctill the waliipc nf Ir- anrl L-Y- / N A i f f e r h x i a fartnr nf rn 2 Thic "('K, 1.- " ,, I" " . "LIII, L l l V . U I U V U "I rb UIlU ,b CR, . 1 1 U l l l V l "J u L U V L W L W I L.U. J . 1 l l l . J difference may reflect differences in the properties of the DP,C and CTAC micelles. The latter are more compact and less ionized than the former. Indeed both the surface area per head group3' and the micelle ionization degree as N increases. Both factors - . L _ 1 L 1 1 * 11 1 r i i 1 1 - r .i 1 0 -nL * --I- I - are expecrea 10 resuit in smaller values or K ana tz ror me excnange or ur; in L I HL than in DP,C micelles as is indeed observed in this work.The values of k+, k- and N listed in table 1 provide a check of the validity of the assumption k+[M] 9 k- which is usually made to obtain [MI and, thus N , when these quantities have not been determined using an immobile quencher. The calculations show that the above approximation would result in an error of 30% at C = 0.055 mol dmP3 and ca. 5% at C 2 0.47 mol dm-3. This is illustrated in fig. 1 by the data points corresponding to DP;. Clearly, the above approximation may introduce large errors at low surfactant concentration and should be avoided whenever possible. Finally, we note that the values of K are fairly inaccurate. This results from the fact that for the present system the condition k- 4 k+[M] is nearly fulfilled. Then K is calculated as the difference between two nearly equal and large numbers, with a large error.(c) Hydrophilic Ionic Qwncher (I-) Eqn (5) has been used to obtain the values of k , and A , - k , which are listed in table 1. The quantity A , - k , which represents the rate of intermicellar exchange of I- [see eqn (1 2)] is plotted against [MI in fig. 2. This rate is seen to increase rapidly and linearly t No value of k,, is given for dodecylpyridinium chloride; however, from the given value of k& for dodecylpyridinium iodide at 25 "C (k& = 3 x lo9 dm3 mol-' s-l) and the relation k,,/k& = c.m.c., with c.m.c. = 1.5 x lo-* mol dm-3, k,, is found to be 4.5 x 10: s-l.116 Fluorescence Quenching in CTAC Micelles 8 7 6 5 v) 3 h W Y I 3 4 2 1 0 I I I I [MI/ 1 0-3 mol dm-3 Fig.2. Variation of A , - KQ = k- (or k,[M] + k-) with micellar concentration. Quenching by A, DP; and +, I-. Dotted line shows the predictions of the ‘hopping’ mechanism. with [MI at low [MI and the plot shows an upward curvature at [MI > 2 x mol dm-3. A linear increase of A , - k , is expected in the whole range of [MI on the basis of the ‘hopping’ mechanisms [see eqn (1 2)], contrary to the present finding. On the other hand, model calculationsz0* 21 predict a strong upward curvature in the change of A , - k , with [MI, even at low [MI, contrary to the results of fig. 2, as well as the results reported by others for the quenching of the fluorescence of pyrenell- l2 and 1-methylpyrene lo by divalent ions in sodium dodecylsulphate micellar solutions, and also for the quenching of pyrene by I- in dodecyltrimethylammonium chloride (DTAC) micellar solutions. Thus the results of fig.2 do not permit a choice between the ‘hopping’ model, and ‘electrostatic’ model. Indeed, both the linear part and the curved part of fig. 2 may be attributed to factors which were not taken into the two models, such as micelle size and shape variations, polydispersi ty, influence of the added ionic quencher, etc. An element of choice between the two mechanisms can however be extracted from the results of fig. 2 when considering the rate of increase of A , - k , with [MI. In the hopping model this rate corresponds to the second order rate constant k , for micelle close encounter with hopping of I- from one micelle to another [see eqn (12)].The plot of fig. 2 yields k, = (1 kO.1) x lo9 dmP3 mol-l s-l at low [MI, and a value about twice as large at high [MI. Our value at low [MI is in excellent agreement with that reported for the pyrene/I- system in DTAC mi~e1les.l~ The rate of diffusion-controlled micellar collisions in the absence of intermicellar interactions is ca. 7 x lo9 dm3 mol-1 s-l at 298.15 K. The strong intermicellar repulsions must significantly decrease this rate, well below the experimental value of lo9 dm3 mol-1 ~ - ~ . t ~ ~ The fact that quencher hopping does not occur at every collision has the same decreasing effect on this rate. Thus, the interpretation of the increase of A , - k , with [MI in terms of micelle close encounters leads to the physically unsound conclusion that the rate for such encounters with I- t This is based on the use of the Debye-Schmoluchowski equation for the rate of diffusion-controlled collisions3@ using for the reaction distance the centre-to-centre intermicellar distance (6 nm) at C = 1 mol dm-‘.The calculations show that if the micelles are assumed to behave like point-charges of charge 6, the rate of collision is already 10 times lower than for electrically neutral micelles. A micelle charge of 8 would result in a reduction of rate by a factor close to 100. Note that the charge of CTAC micelles is ca. 25.A . Malliaris, J. Lang and R. Zana 117 hopping is larger than the rate of diffusion-controlled micellar collisions. This inter- pretation must therefore be discarded and we are led to the conclusion that the increase of A , - k , with [MI probably has its origin in the electrostatic effects postulated by Almgren et al.20,21 Whichever the model adopted for the quencher exchange, the extrapolation of the plot of fig.2 to [MI = 0 yields the rate constant k; for the dissociation of I- from CTAC micelles, in the absence of intermicellar interactions [see eqn (12)]. The value of k; = 2.5 lo5 s-l obtained in the present study is one order of magnitude smaller than that found by Grieser13 for the pyrene/I-system in DTAC (a shorter homologue of CTAC) micelles. This large difference may reflect the smaller surface area per head group of the CTAC micelles with respect to DTAC micelles, which results in a tighter binding of counterions including I-, and thus in a smaller k- value, in the absence of intermicellar interactions. Attempts made to calculate k+ from the results produced largely scattered values, some even negative.This effect results from the fact that the approximation k+[W B k- is nearly fulfilled at all [MI in the case of I- quenching, and thus, k+ cannot be obtained by the fluorescence method with any accuracy. The fact that this approximation is fulfilled explains that the two sets of data in columns (c) and (6) in table 1 are nearly coincident. The value of k , for quenching by iodide is slightly larger in CTAC (this work) than in DTACL3 micelles. The difference probably reflects the fact that the latter study was performed at 22 "C. The k, values for I- quenching show the same decreasing trend upon increasing N than for CP; or DP; quenching, and the product Nk, is also nearly constant.At a given N the k , value for pyrene is smaller than for the other three quenchers. This probably reflects the well known interaction between arenes, such as p~rene,~O* 41 and the trimethylammonium head group, which restricts the mobility of pyrene. This interaction is also reflected in the much lower k , values found for all cationic micelles than for anionic micelles such as SDS (k, z 2 x 10' s-l),'? in the case of pyrene quenching . Conclusions This study has permitted us to show that CTAC micelles remain nearly spherical up to a concentration of 1.2 mol dm-3, i.e. at concentrations very close to those for the formation of liquid crystal phases. lu'evertheless, the aggregation number is increased by a factor of two in going from the c.m.c.to 1.2 mol dmP3. Thus, the assumption sometimes made13 in similar studies that N remains constant over large C range is not valid. We have also shown that great care should be exercised in the approximations made for the evaluation of N by means of transient fluorescence quenching methods. Finally, our results suggest that the change of the rate of dissociation of I- from the CTAC micelles upon increasing micelle concentration is due to the electrostatic effect first discussed by Almgren et al. We thank the PIRSEM (CNRS) for its financial support (AIP No. 2004). A.M. thanks the French Ministere de l'lndustrie et de la Recherche for financing his stay in Strasbourg. References 1 S. S. Atik and L.A. Singer, Chem. Phys. Lett., 1978, 59, 519. 2 M. Almgren and J. E. Lefroth, J . Colloid Interface Sci., 1981, 81. 486. 3 S. S . Atik and J. K. Thomas, J . Am. Chem. Soc., 1981, 103, 3550. 4 S. S. Atik and J. K . Thomas, J . Am. Chem. SOC., 1981, 103, 4367. 5 J. E. Lofroth and M. Almgren, J . Phys. Chem., 1982, 86, 1636.118 Fluorescence Quenching in CTAC Miceflrs 6 M. Van der Auweraer, C. Dederen, C. Palmans-Windels and F. C. De Schryver, J. Am. Chem. Soc., 7 E. Roelants, E. Gelade, M. Van der Auweiaer, Y . Croonen and F. C. De Schryver, J . Colloid Interface 8 Y. Croonen, E. Gelade, M. Van der Zegel, M. Van der Auweraer, H. Vandendriessche, F. C. De 9 M. A. J. Rodgers, E. Da Silva and M. F. Wheeler, Chem. Phys. Lett., 1978, 53, 165. 1982, 104, 1800.Sci., 1983, 96, 288. Schryver and M. Almgren, J . Phys. Chem., 1983, 87, 1426. 10 J. C. Dederen, M. Van der Auweraer and F. C. De Schryver, Chem. Phys. Lett., 1979, 68, 451. 11 F. Grieser and R. Tausch-Treml, J . Am. Chem. Soc., 1980, 102, 7258. 12 J. C. Dederen, M. Van der Auweraer and F. C. De Schryver. J. Phys. Chem., 1981, 85, 1198. 13 F. Grieser, Chem. Phys. Lett., 1981, 83, 59. 14 S. S. Atik, M. Nam and L. Singer, Chem. Phys. Lett., 1979, 67. 75. 15 P. Lianos, J. Lang and R. Zana, J. Phys. Chem., 1984, 88, 819. 16 S. S. Atik and J. K. Thomas, J . Am. Chem. SOC., 1982, 104, 5868. 17 P. Tnfelta, Chem. Phys. Lett., 1979, 61, 88. 18 A. Yekta, M. Aikawa and N. J. Turro, Chem. Phys. Lett., 1979, 63, 543. 19 A. Henglein and Th. Proske, Ber. Bunsenges. Phys. Chem., 1978,82,471; Y. Moroi, A. M. Braun and 20 M. Almgren, G. Gunnardson and P. Linse, Chem. Phys. Lett., 1982, 85, 451. 21 M. Almgren, P. Linse, M. Van der Auweraer, F. C. De Schryver, E. Gelade and Y. Croonen, J . Phvs. 22 A. Malliaris, J. Lang and R. Zana, J . Colloid Interface Sci., in press. 23 P. Infelta and M. Gratzel, J . Phys. Chem., 1979, 70, 179. 24 G. Pfeffer, H. Lami, G. Laustnat and A. Coche, C.R. Acad. Sci., Paris, 1963, 257. 434. 25 P. Infelta, M. Gratzel and J. K. Thomas, J. Phys. Chem., 1974, 78, 190. 26 M. Tachiya, Chem. Phys. Lett., 1975, 33, 289. 27 M. Gueron and G. Weisbuch, Biopolymers, 1980, 19, 353. 28 B. Lindman and H. Wennerstrom, Top. Curr. Chem., 1980, 87, 1. 29 G. Manning, Annu. Rev. Phys. Chem., 1972, 23, 117. 30 C. L. Kwan, S. S. Atik and L. A. Singer, J . Am. Chem. Soc., 1978, 100, 4783. 31 H. V. Tartar, J . Phys. Chem., 1955, 59, 1195; C. Tanford, J. Phys. Chem., 1972, 76, 3020. 32 M. Van der Auweraer, J. C . Dederen, E. GCladC and F. C. De Schryver, J . Chem. Phys., 1981,74, i 140. 33 F. Reiss-Husson and V. Luzzati, J. Phys. Chem., 1964,68. 3904. 34 G. Porte, Y. Pogg, J. Appell and G. Maret, J. Phys. Chem., 1984, 88, 5713. 35 B. Ninham, D. Evans and G. Wei, J . Phys. Chem., 1983, 87, 5020. 36 H. Hoffmann, R. Nagel, G. Platz and W. Ulbricht, Colloid Polymer Sci., 1976, 254, 812. 37 E. A. G. Aniansson and S. N. Wall, J. Phys. Chem., 1974, 78, 1024; 1975, 79, 857. 38 R. Zana, J. ColloidInterface Sci., 1980, 78, 330. 39 M. Eiger and L. De Maeyer in Techniques of Organic Chemistry, ed. S. Friess, E. Lewis and A. Weissberger (Interscience, New York, 1963), vol. VIII, Part 11, p. 1032. 40 P. Lianos, M-L. Viriot and R. Zana, J. Phys. Chem., 1984, 88, 1098. 41 M. Almgren, F. Grieser and J. K. Thomas, J . Am. Chem. SOC., 1979, 101, 279. M. Gratzel, J . Am. Chem. SOC., 1979, 101, 567. Chem., 1984, 88, 289. Paper 51494; Receizjed 25th March, 1985

ISSN:0300-9599    

DOI:10.1039/F19868200109

出版商:RSC    

年代:1986

数据来源: RSC

摘要:

J. Chem. SOC., Faruday Trans. 1, 1986, 82, 119-123 The Metachromism of Methyl Orange with Electrolytes and Possible Salting-out J. Graham Dawber," David T. Fisher and Paul R. Warhurst Department of Chemistry and Biology, North Stafordshire Polytechnic, Stoke-on-Trent ST4 2DE The changes in optical absorption of the Methyl Orange (MO) anion in solution produced when electrolytes are added has been studied and confirms the work of others. However, the measurement of circular dichroism and also absorption spectra on filtered and unfiltered solutions of MO in the presence of high concentrations of electrolyte suggest that a new absorption band at 360 nm may be related to precipitation of the MO. ~~ _________ -~ p- Three recent paperslp3 have drawn attention to the metachromic effect of the addition of electrolytes to basic solutions (pH 8-10) of Methyl Orange (MO).The results have been interpreted in terms of the solvent structure-breaking properties of the electrolytes and the subsequent effect upon the competitive solvation of the Methyl Orange anion (I) and also ion-pair formation. We had considered the possibility of stacking of the MO as being a possible contributing factor to the metachromic behaviour, and this led us to measure the circular dichroism (c.d.) of the solutions of MO in electrolyte solutions. Experimental A stock solution of Methyl Orange (B.D.H.; used as received) of 4.28 x mol dmP3 was prepared and the pH adjusted by the addition of dilute NaOH. The electrolytes studied included KCI, KI, MgCI, and CaCl,, all of which were of AnalaR quality.The stock MO solution was diluted tenfold by solutions of the electrolytes in deionised water so that the MO solution for most measurements was 4.28 x 10-5 mol dmp3 and had a The absorption spectra of the solutions were measured in silica cells using a Varian DMS 90 spectrophotometer. The spectrum of each MO solution was measured relative to a reference solution containing the same concentration of electrolyte. A,,, for MO in water at pH 10 was found to be 463 nm, and the molar absorbance, E , at this wavelength was 2.33 x lo5 dm3 mol-l cm-l. The c.d. measurements were carried out using a home-built single-beam instrument based upon the monochromator of a Hilger Uvispek spectrophotometer which has been fitted with a motorised wavelength drive. Light from the monochromator is linearly polarised by a Rochon prism and then modulated at 50 kHz into right- and left-circularly polarised light by a Morvue photoelastic modulator.The light beam after passing through the sample falls on a photomultiplier (EM1 9789Q/B) the output from which is fed to a simple circuit to separate the a.c. component (arising from the c.d. of the sample) from the d.c. component. The a.c. component is fed to a Brookdeal 9501 lock-in amplifier and the processed signal is recorded on a potentiometric chart recorder. While the c.d. spectrum is being recorded the d.c. signal from the photomultiplier is maintained pH of 9-10. 119120 Metachromism of Methyl Orange 450 400 3 50 h/nm Fig. 1. Circular dichroism of MO (8.6 x mol dmP3) in 2 mol dm-3 KC1: (a) unfiltered solution; (h) filtered solution.at a constant level by adjusting the e.h.t. to the photomultiplier. This procedure then compensates the measurements for changes in light-source output and photomultiplier sensitivity with wavelength of light. The c.d. spectrometer was calibrated with an aqueous solution of D-camphor sulphonic acid4 and also a solution of epiandrosterone in dioxan. Results and Discussion For all the electrolytes studied the results confirmed the findings of de Vijlder,1-3 namely that the optical absorbance of the Methyl Orange anion is decreased by the presence of the electrolyte. However, in a number of our experiments with electrolyte at high concentration, and particularly with MgCl,, which is very effective in lowering the absorbance of the MO, it was found that when the glassware was emptied and then rinsed with a small quantity of acetone, the acetone became orange in colour and its absorbance was appreciable at the wavelength corresponding to that of Amax of MO.This suggested that MO might be precipitated at high electrolyte concentration and adsorbed onto the glassware in our experiments. In addition, some solutions at high electrolyte concentration exhibited a slight turbidity, a feature which had previously been noted.lI A number of planar organic molecules have been shown to form face-to-face dimers and even higher aggregates in aqueous solution, often arising from hydrophobic interactions between the molecules in which water is excluded. Examples include phena~ine,~ tacrine hydrochloride,6 Acridine Orange,i Methylene Blue* and thi~nine.~ Admittedly all of these compounds are cationic, but in principle there should be no reason why planar molecules such as the Methyl Orange anion should not stack in a similar way. If higher aggregates were formed in solution by stacking of the MO, then theJ. G.Dawber, D. T. Fisher and P . R. Warhurst 121 100 -50 1 I I I L 1 I 1 420 L 00 380 360 A/ nm Fig. 2. Circular dichroism of MO (4.28 x lop4 mol dm-3) in 4 mol dm- KI: (a) unfiltered solution; (b) filtered solution. N.b. the larger c.d. with a more concentrated MO solution and in which suspended material was more evident. possibility exists of observing c.d. arising from a stacked, inherently symmetric chromo- phore.It was this possibility which prompted us to measure the c.d. of MO solutions in the presence of electrolyte within a wavelength region where a new absorption band has been reported.l$ The MO itself in solution did not exhibit any c.d., but in concentrated solutions of KCl and KI an appreciable c.d. was observed within the normal absorption band of MO at 460 nm, but also in the wavelength band region of the new absorption band at ca. 360 nm.l9 In the latter case the c.d. was biphasic. On very close examination of the solutions they were seen to contain traces of almost transparent microcrystalline platelets. Since c.d. measurements are very sensitive to the presence of suspensions, the solutions were filtered through a fine filter paper (Whatman no. 542) and the c.d.remeasured. The c.d. was found to have disappeared over the whole wavelength range (fig. 1 and 2), and furthermore there were traces of MO left on the filter paper. The presence of precipitated MO was quite apparent when the solutions containing high electrolyte concentrations were filtered by drawing them into a Pasteur dropper with a small piece of paper tissue wrapped round its end. A similar effect was observed with MO in 1.7 mol dmP3 CaCl, solution when c.d. was observed within the new absorption band at 360 nm, which then disappeared on filtration of the solution (fig. 3). Thus the c.d. observed was not due to stacking of the MO, but somehow related to the presence of the microcrystalline platelets which had precipitated in the concentrated electrolyte solutions.A further piece of evidence suggesting that salting-out of the MO as being related to the new absorption band came from the absorption spectra themselves. Fig. 4 shows the spectrum of MO in water (diluted to fall in a lower absorbance range) and spectra of unfiltered and filtered MO in 1.7 mol dm-3 CaCl,. The unfiltered solution of MO in the122 Metachromism of Methyl Orange 4 00 380 360 3LO X/ nm solution; (b) filtered solution. Fig. 3. Circular dichroism of MO (8.6 x lop5 mol dm-3) in 1.7 mol dm-3 CaC1, : (a) unfiltered X/ nm Fig. 4. Absorbance spectra of MO in 1.7 mol dmP3 CaCI, solution: (a) in water (solution diluted to fit on same scale); (6) in 1.7 mol dm-3 CaCl,, unfiltered solution ([MO] = 4.28 x lop5 mol dmP3); (c) in 1.7 rnol dm-3 CaCl,, filtered solution ([MO] = 4.28 x mol dm-”).J .G. Dawber, D. T. Fisher and P. R. Warhurst 123 CaCl, solution shows an absorption band at ca. 360 nm (the new absorption band1). The corresponding filtered solution, however, shows no such band at 360 nm, although the main absorption at 460 nm is slightly reduced (possibly owing to adsorption or removal of MO on the filter paper). This implies that the appearance of an absorption band at 360 nm is related in some way to the suspended material produced by the high concentration of electrolyte. The financial assistance of the S.E.R.C. towards the construction of the c.d. spectrometer is grate fully acknowledged. References 1 M. De Vijlder, J . Chem. Soc., Faraday Trans. I , 1982, 78, 137. 2 M. De Vijlder, J . Chem. Soc., Faraday Trans. I , 1983, 79, 155. 3 M. De Vijlder, J . Chem. Soc., Faraday Trans. I , 1985, 81, 1369. 4 J. Y. Cassim and J. T. Yang, Biochemistry, 1969, 8, 1947. 5 D. Attwood, J . Chem. SOC., Faraday Trans. I , 1983, 79, 2669. 6 J. Gormally, N. Natarajan, E. Wyn-Jones, D.Attwood, J. Gibson and D. G. Hall, J . Chem. SOC., 7 R. Cohen and S. Yariv, J . Chem. Soc., Faraday Trans. I , 1984, 80, 1705. 8 S. L. Fornili, G. Sgroi and V. Izzo, J . Chem. SOC., Faradq Trans. I , 1983, 79. 1085. 9 S. L. Fornili, G. Sgroi, L. Palumbo and V. Izzo, J . Chem. Soc., Faraday Trans. I , 1985, 81, 255. Faradq Trans. 2, 1984, 80, 243. Paper 51520; Receiijed 27th Murch, 1985

ISSN:0300-9599    

DOI:10.1039/F19868200119

出版商:RSC    

年代:1986

数据来源: RSC

摘要:

J. Chem. SOC., Furuduy Trans. I , 1986, 82, 125-142 Interfacial Tension Minima in Oil-Water-Surfactant Systems Behaviour of Alkane-Aqueous NaCl Systems containing Aerosol OT Robert Aveyard," Bernard P. Binks, Steven Clark and Jeremy Mead Chemistry Department, University of Hull, Hull HU6 7RX In heptane and aqueous NaCl systems containing diethylhexyl sodium sulphosuccinate (AOT) very low interfacial tensions, y, can be attained. For fixed temperature (T) and salt concentration, y falls as the AOT concentra- tion increases and levels off at a value of yc at a concentration (the critical micelle concentration, c.m.c.) corresponding to the onset of surfactant aggregation in either the aqueous or oil phase, depending on T and salt concentration. For a given T, at low salt concentration AOT resides in the aqueous phase both below and above the c.m.c. At higher salt concentrations surfactant transfers to the oil phase and leaves the aqueous phase close to the c.m.c.expected if no excess of oil were present, but devoid of micelles. In this case the oil phase is shown to be a dilute water-in-oil microemulsion. The transition between behaviour at high and low salt concentrations corresponds to the attainment of a minimum yc. The variation in yr with salt concentration is in part a consequence of the way in which the c.m.c. and the surfactant activity coefficients change with salt concentration. The minimum yc occurs when the apparent degree of dissociation of surfactant in the micelle and at the oil-water interface are equal and close to zero. -- _ ~ _ _ _ _ _ _ ~ There is a continuing interest in the occurrence of ultra-low interfacial tensions in oil-water systems containing either ionic or non-ionic surfactants.Tensions in such systems are of intrinsic interest but are also relevant to the understanding of various phenomena, e.g. the thermodynamics of formation of microemulsions,l phase inversion of macroemu1sions2 and the process of enhanced oil recovery by surfactant flooding of oil wells.3 A limited amount of work has been reported on systems containing pure surfactant (as opposed to commercial mixtures). Recently Franses and coworkers4-6 have reported studies of systems containing 8-phenylhexadecane sodium sulphonate (Texas I) which forms liquid crystallites rather than micelles in aqueous solution, at least under the conditions studied.Alkane-aqueous NaCl systems can yield ultra-low tensions in the presence of Texas I, but these tensions are not reproducible, as a result of the formation of a liquid-crystalline layer at the oil-water interface. Commercial surfactant systems have been studied widely [see, for example, ref. (3) and (7)]. Chan and Shah3 report an investigation into the behaviour of commercial petroleum sulphonates in oil-water systems. They conclude that in such systems a minimum tension is attained when the equilibrium aqueous phase is at its critical micelle concentration (c.m.c.) and simultaneously the distribution coefficient of surfactant between oil and water is unity. It is said that this condition corresponds to maximum surface excess of surfactant in the monolayer adsorbed at the oil-water interface.The aim of the present work has been to ascertain in detail how a pure anionic surfactant (diethylhexyl sodium sulphosuccinate, AOT) which is known to form micelles in dilute aqueous solutions is distributed between aqueous solution and a normal alkane (heptane), and how the distribution depends upon [surfactant], [salt] and temperature, T. 125126 Tension Minima in Heptune-Aqueous NaCl-A0 T Systems The distribution behaviour is related to the occurrence of interfacial tension minima, and the tension data are analysed using a recently presented thermodynamic approach. The structure of equilibrium alkane phases which contain surfactant and water have also been investigated. Experimental Materials The heptane was a puriss sample with a purity > 99% as confirmed by g.1.c.in this laboratory, and it was passed through chromatographic alumina before use. Water was distilled once, passed through an Elgastat ion-exchange column and then through a Milli-Q-Reagent water system. The heptane-water interfacial tension at 25 "C (determined by the Wilhelmy-plate method) was 50.80 mN m-l, in good agreement with a previous value of 50.77 mN m-l.g Sodium chloride was AnalaR grade and was used untreated after confirming that aqueous solutions had surface tensions in good agreement with those previously obtained using samples roasted for 12 h at 450 OC.l0 The AOT was from Sigma; two different batches were employed and gave complete mutual agreement in the distribution and interfacial-tension experiments.The purity, as determined by the hyamine titration method,ll was 99.5 % . No minima were observed in plots of interfacial tension against In [surfactant] (see later); such plots (fig. 9) were of the normal shape for a micelle-forming surfactant. In our experiments we used the AOT in the presence of a swamping excess of NaC1, and we do not expect small amounts of uni-univalent electrolyte impurities to sensibly affect results. However, the presence of divalent metal ions could significantly affect the surfactant behaviour.12 We found, by atomic absorption spectrometry, that the AOT contained 0.5 &- 0.2 mol% Ca2+ ions. However, in the worst case the ratio Ca2+: Na+ was 1 : 7000, and we have ascertained that addition of Ca2+ at the levels already present has no detectable affect on the low oil-water interfacial tensions of interest.Methods The distribution of surfactant between heptane and aqueous phases at equilibrium was determined with respect to both salt concentration and surfactant concentration. Typically, heptane and aqueous solutions of NaCl were shaken with the surfactant initially in either phase, and the resulting emulsion was left to separate for a few days in a thermostat; the two phases were then separated. The surfactant concentrations were determined by titration with hyamine.ll The heptane phases were also analysed for water by the Karl-Fischer method13 using a Baird and Tatlock AF3 automatic titrator. Viscosities of water-in-oil microemulsions were determined using Ubbelohde visco- meters which were calibrated by measuring flow times of various pure n-alkanes and water-glycerol mixtures.For the measurements a viscometer was suspended vertically in a water bath thermostatted at 25.00+0.05 "C. The densities of the liquids were determined using a Paar DMA 55 densimeter. Emulsion conductances were determined using a Jenway PCM3 digital conductivity meter which had a facility for the digital read-out of temperature. Interfacial tensions above ca. 3 mN m-l were measured using a Kruss K10 automatic tensiometer. For tensions above ca. 10 mN m-' both the du Nouy ring and Wilhelmy plate methods were employed, giving results in good mutual agreement (i.e. within 0.1 mN m-l). Below 10 mN m-l, however, the plate method became unsatisfactory and only the ring method could be used.Tensions below 3 mN m-l were determined using a Kruss spinning-drop tensiometer. Cu. 1 mm3, depending on the tension, of the oil phase was injected into the horizontalR. Aueyard, B. P. Binks, S. Clark and J. Mead 1 E Y .C Y Y 0 u 2- T/"C [ NaCl] /mol dmW3 (a) Conductivity of stirred emulsions as a function of temperature for aqueous 127 NaCl concentrations: 0,0.068; @, 0.086; .,0.103 mol dm-3. Aqueous phase volume fraction = 0.40. (b) Conductivity of stirred emulsions as a function of salt concentration at 25 "C for aqueous-phase volume fractions of 0.40 (0) and 0.67 (O), respectively (@ are overlaying points). All systems contained 5 wt % AOT overall. capillary (rotating about its long axis) filled with the aqueous phase.At sufficiently high frequencies the drop assumes the form of a cylinder of radius r with hemispherical ends, and the interfacial tension, y, is then given by y = r3Apn2n2 (1) where n is the frequency of rotation and Ap the difference in density between the two phases. Tensions reached an equilibrium value in a few minutes (see later). Photon correlation spectrometry (P.c.s.) was performed using a Malvern Instruments PCS 100 spectrometer. Polarised light from a 15 mW He-Ne laser was focused on the sample in a 1 cm fluorescence cuvette in a thermostatted container. Scattered light was detected at 90" to the incident beam and the correlation function evaluated using a Malvern K7027 correlator. Correlation functions were analysed by the method of cumulants giving a mean correlation length and a variance of the distribution of e~ponentia1s.l~ In the limit of infinite dilution of a microemulsion the correlation length can be equated to the hydrodynamic radius of the scattering particles.Results and Discussion We will be interested in the effects of surfactant concentration, salt concentration and temperature on y ; effects of alkane chain length will be discussed e1~ewhere.l~ All these variables affect the distribution of surfactants between aqueous and oil phases and the type of emulsion formed on agitating the systems.16 Emulsion Type The emulsion type (oil-in-water, O/W, or water-in-oil, W/O) is readily characterised by conductivity. In fig. 1 ( a ) we show how, for constant salt concentration, an increase in Tcan lead to inversion from W/O to O/W emulsions.The former, which have oil as the continuous phase, exhibit low conductivity, whereas for the latter emulsion type the conductivity is high. The phase-inversion region is shifted to higher temperatures as 5 F A R 1128 Tension Minima in Heptane-Aqueous NaCI-A0 T Systems P "- a a I I 1 0 - ~ [ AOTIinitial/~nol dmr3 Fig. 2. Initial and equilibrium aqueous-phase concentrations of AOT in heptane-aqueous NaCl systems at 25 "C. The diagonal line has unit slope. Sodium chloride concentrations are: 0,0.0171; (0, 0.0513; 0 , 0.1027 mol dm-3. 0 5 10 [ AOT] ,nitial / 1 0-3 mol Fig. 3. Initial and equilibrium heptane phase concentrations of AOT in heptane-aqueous NaCl systems at 25 "C. Slope of line through filled points is unity.For a salt concentration of 0.0 17 mol dmP3 no surfactant detectable in heptane. Points (0) and (0) are for NaCi concentrations of 0.1027 and 0.0513 mol dmP3, respectively. The slope of less than unity through the points (0) indicates a loss of surfactant to a third, surfactant-rich phase.R. Aveyard, B. P. Binks, S. Clark and J . Mead 129 100 0 I I b 0 0.05 0.10 [ NaCl],q/mol dmA3 Fig. 4. Distribution of AOT (in aggregated form) between heptane and aqueous NaCl. Points (a) represent the percentage of AOT in the heptane phase, given by [AOTIhept #,,,,/([AOT],, - #aqu c.m.c.), and points (0) represent the percentage of AOT in the aqueous phase, given by ([AOT],,, - c.m.c.) #,,/([AOT],, - #a,u c.m.c.). [AOT],, is the overall surfactant concentration (0.05 mol dm-3) and # are volume fractions.(a) 10, (h) 25 and (c) 40 "C. the salt concentration is increased. For constant temperature, if the emulsion is initially O/W, addition of electrolyte forces an inversion to W/O, as seen in fig. 1 (b). For constant surfactant concentration the salt concentration required to effect inversion depends on the phase volume ratio, and for this reason we will not be able to obtain quantitative correlation between these data and the tensions to be discussed later. Surfactant Distribution Under conditions where W/O emulsions are formed, surfactant is present in the oil phase, and (as will be seen) the oil phase is itself a W/O microemulsion. We have studied surfactant distribution between aqueous and alkane phases in some detail. For low salt concentrations, as the surfactant concentration in the system is increased surfactant resides totally in the aqueous phase, which can be either above or below the c.m.c., as depicted in fig.2. At higher salt concentrations, however, although the AOT remains totally in the aqueous phase up to a concentration approximately equal to the c.m.c. expected in an aqueous phase saturated with alkane (but with no excess alkane phase present), at concentrations greater than this the surfactant is also present in the alkane phase, the aqueous phase concentration remaining constant (fig. 3). Under these conditions no micelles are detectable (by P.c.s.) in the aqueous phase. The determination of critical micelle concentrations and their significance are discussed in the Appendix.5-2130 a o 70 R 60 50 - I I 0 - ( a 1 I I In fig. 4 the effect of salt concentration on distribution at three temperatures is shown; the overall AOT concentration in the system is the same for all the experiments. At low salt concentrations the surfactant is all in the aqueous phase, which is above the c.m.c. As the salt concentration is increased transfer of surfactant to the alkane occurs and the aqueous phase is left, devoid of micelles, close to its expected c.m.c. At higher temperatures, higher salt concentrations are needed to effect surfactant transfer to oil, in accord with the conductivity data. At intermediate salt concentrations, close to conditions corresponding to phase inversion, a third surfactant-rich phase is formed which contains much of the surfactant in the system, as can be appreciated from fig.3. The formation of such third phases in oil-water-surfactant systems is well documented.l79 l8 Nature of the Equilibrium Oil Phases Under conditions where surfactant is detected in the heptane, water is also present. For constant salt concentration and temperature the ratio (mol water)/(mol AOT) in heptane ( R ) is found, within experimental error, to be independent of oil-phase surfactant concentration, as seen in fig. 5(a). On the other hand R depends on salt concentration [fig. 5(h)]. As the salt concentration increases R decreases.R . Aueyard, B. P . Binks, S . Clark and J . Mead 131 0 0.1 0.2 C/g ~ 3 7 1 ~ ~ Fig. 6. Viscosities of W/O microemulsions in the system AOT-H,O-heptane at 25 "C for a range of values of droplet concentration C and of R.Points are: 0, R = 20; 0, R = 30; @, R = 40; @, R = 50; 0, R = 60. Since surfactant is not detected in the oil phase below the c.m.c., we may suppose that when surfactant is present it is in aggregated form. Further, since substantial amounts of water are only found in conjunction with surfactant, we assume the water is in the form of droplets coated with surfactant monolayers. The existence of such droplets is demonstrated by P.C.S. measurements. Information on the shape and/or solvation of the droplets can be gleaned from viscosity data. The specific viscosity, qsp, of a solution (microemulsion here) is defined as (q -qsolv)/ qsolv, where q and qsvlv are the viscosities of the solution and the solvent (heptane). In fig.6 are shown plots of qsp/C against C (the droplet concentration in g cmP3 of solution) for five different values of R at 25 "C. The data have been fitted to quadratic (2) equations of the form where a and b are constants and [q] is the intrinsic viscosity. A knowledge of [q] is useful, since it is relatively insensitive to the size of any particles (aggregates), but can give information on the shape and/or solvation of the particles. For the present systems [q] qsp/C = [q] + aC+ bCz may be expressed as19 [q] = v( VD + SV,) (3) where VD and V s are partial specific volumes of the droplets and heptane, respectively, and 6 is the mass of heptane per unit mass of droplet. The quantity v is the Simha factorz0 and is 2.5 for spherical particles.We assume that the density, ps, of heptane is unchanged132 Tension Minima in Heptune-Aqueous NaC1-A0 T Systems Table 1. Solvation and Simha factors for W/O microemulsions in heptane stabilised by AOT at 25 "C ~- ~~~ wt To heptane R [rl v,(6 = 0) ( V = 2.5) 20 3.0 1 2.93 13.9 30 2.85 2.79 9.5 40 3.1 1 3.05 16.5 50 3.03 2.98 14.3 60 3.28 3.23 21.1 ~~ ~~~ ~~ ~~ ~~ ~~ in microemulsions, so that 6 = pgl = 1.472 cm3 g-'. Values of vD at each value of R were obtained from the measured densities, pM, of the microemulsions, noting that where WD and W, are the weight fractions of droplets and of heptane. Values of v obtained from eqn (3) by assuming 6 = 0 (no solvation), together with the wt % of heptane associated with the droplets on the assumption of spherical aggregates (i.e.setting v = 2.5) are given in table 1. Assuming 6 = 0, the values of v obtained indicate axial ratios of 1 .8-2.8.21 Intuitively, however, we feel a more acceptable interpretation of the viscosity data is that the droplets are essentially spherical but solvated by heptane. If so, the degree of solvation amounts to between ca. 10 and 20%, which is approximately one heptane molecule associated with every two surfactant molecules at the droplet interface. Penetration of alkane into the surfactant chain region around the droplets is to be expected, as discussed by Mukherjee et a/.22 A reasonable estimate of the droplet size can be made on the assumption that the droplets are monodisperse spheres. From simple geometry the radius, rC, of the water core is where AD is the area available per surfactant ion at the drop surface, and u, is the volume occupied by a water molecule in liquid water (0.03 nm3).The hydrodynamic radius, rh, discussed later is expected to exceed r , by approximately the length of the surfactant ions (ca. 1 nm). If we take A, to be equal to the area, A,. occupied in a close-packed monolayer at the plane heptane-aqueous NaC1 interface (0.73 nm2, see later), and R as 67, [fig. 5(a)], r , = 8.3 nm, so that rh is calculated to be ca. 9.3 nm; values of Yh are referred to again later. In any event it is apparent that those equilibrium oil phases which contain water are dilute W/O microemulsions. Values of Yh can be obtained from P.C.S. measurements, and W/O microemulsions stabilised by AOT in heptane have previously been studied both by P .C . S . ~ ~ and by small-angle neutron scattering (SANS).24 The work we report extends the droplet size range studied by Nicholson et a/.23 Unlike previous workers, we have investigated microemulsions in equilibrium with excess phases [the system represented in fig. 5(6)] as well as 'made-up' samples prepared by addition of the required amounts of water and surfxtant to the pure hydrocarbon. No salt was included in the made-up samples, and we return to this point later. The measured correlation length, I, obtained in a P.C.S. experiment contains a contribution from interparticle interactions and can only be equated with rh at infinite dilution of droplets. We have determined I over a range of volume fraction 4D of dispersed phase for all the microemulsions studied. Values of IF1 were plotted against dD to obtain infinite-dilution values of 1; sample plots are shown in fig.7. Values of Yh (I at 4D = 0) are plotted against R in fig. 8, where results from ref. (23) are also included. It is clear from eqn (9, if A, is constant, that such a plot should be linear and that A , ( 5 ) rc = 3 Ru,(AD)-'R. Aveyard, B. P. Binks, S. Clark and J . Mead 133 0.100 6 I E -50.075 M I - 0.050 I I 0.1 0.2 4D Fig. 7. Variation of correlation length, 1, with droplet volume fraction, &,, for W/O microemulsions at 25 “C. Points a, 0 and 0 are for R = 46, 60 and 73, respectively. can be obtained from the slope. The intercept at R = 0 gives the effective thickness, t, of the surfactant film.A linear least-squares fit of the (rh, R ) data gives t = 1.4 0.1 nm and A , = 0.50+0.03 nm2. The results for the made-up and equilibrium microemulsions fall on a common line and yield a value of A , which requires comment. The value of A , of 0.73 nm2 mentioned earlier refers to a system containing a swamping excess of NaCl (see next section), whereas the made-up microemulsions contained no NaC1. However, the counterion (Na+) concentration in microemulsion droplets is high (ca. 2 mol dm-3 for rc = 5 nm and A , = 0.50 nm2) and presumably constitutes swamping excess of ‘salt’, so A , and A , should not differ on this count. The difference could arise as a result of the molecular geometry of AOT. At a plane interface surfactant ions pack more closely as the salt concentration is increased, until the effective headgroup cross-sectional area (determined by electrical repulsion and solvation) becomes equal to that of the hydrocarbon tail region.It is possible, however, that the headgroup area may be reduced still further until its ‘geometrical’ size (ca. 0.45 nm2) is reached. At a plane interface A , cannot be reduced below the tail area, but at the convex surface of a water droplet the area per molecule can fall below A,. There are, however, other possible origins for the differences between A , and A,. One is that there may be a significant amount of AOT dissolved in the heptane rather than adsorbed at the droplet surfaces. We feel, however, that this is unlikely on the basis of our distribution measurements.A more likely alternative is a degree of polydispersity of drop size.23134 Tension Minima in Heptane-Aqueous NaC1-A0 T Systems 20 E 0 50 R 100 Fig. 8. Dependence of hydrodynamic radius, rh, on R at 25 "C. Points (0) and (9) are for 'made-up' and equilibrium systems, respectively, in the present study; 0, ref. (23); a, values of rC [obtained from SANS, ref. (24)] with 1.2 nm added (see text). 20 - I E z E 10 3 0 -11 -8 -5 In [ AOT]/mol d111-~) Fig. 9. Interfacial tensions between heptane and aqueous phase at 25 "C as a function of AOT concentration in the aqueous phase. Aqueous-phase NaCl concentrations are: a, 0.0256; 0 , 0.0513; 0, 0.1027 mol drnp3.R. Aveyard, B. P. Binks, S. Clark and J. Mead 135 0 0.05 0.10 rnN,/mol dme3 Fig. 10. Variation of yc with salt concentration at 25 "C.In all cases the AOT concentration is above the c.m.c., so aggregates are present in the heptane or the aqueous phase. Points (0) are for non-pre-equilibrated phases and (0) are for equilibrium phases (see text). The full line is obtained using eqn (8) with B = 1.078 mN m-l. Interfacial tensions We denote the surfactant concentration by mA. The shapes of the plots of the interfacial tension against lnm, are those expected for an ionic, micelle-forming surfactant in the presence of added electrolyte (fig. 9). Below the break point, which corresponds to the c.m.c., the plots are h e a r over a wide range of mA. In this region the surface excess of surfactant, rA (obtained using the appropriate form of the Gibbs equation)26 is constant and corresponds to the presence of a saturated monolayer.The close-packed area, A , (= 1 /FA), for salt concentrations of 0.05 mol dmP3 and higher is ca. 0.73 nm2, as already mentioned. Above the c.m.c. y remains constant at 7,. As seen earlier, aggregation of surfactant occurs in either the aqueous or the oil phase (but not both simultaneously), depending on conditions. The c.m.c. as defined here is the aqueous-phase concentration of surfactant which corresponds to the onset of aggregation in the preferred phase. In a certain range of salt concentration the values of y c become very low (1 0-3 to lop2 mN m-l) and, from the shape of the plots of y against In mA we conclude that such low tensions can be obtained through monolayer adsorption. In experiments leading to the y , against salt concentration curves and the yc against temperature curves, a drop of pure heptane is introduced into the capillary of the spinning-drop tensiometer, which is filled with aqueous surfactant above the c.m.c.The tension recorded is found to be independent of surfactant concentration (above the c.m.c.) and is thus taken to be equal to yc. That this is entirely reasonable can be seen from the results shown in fig. 10, where four different sets of data are represented. First136 Tension Minima in Heptane-Aqueous NaC1-A0 T Systems r 1 0 0.05 0.10 [ NaCl] /mol dm-3 Fig. 11. Correlation between surfactant transfer between phases and minimum yc at 25 "C. 0 , Tensions; 0, percentage aggregated AOT in heptane phase. the points obtained as just described are shown, Also included are the yc values from the y against In mA plots (fig.9), and those for 25 "C taken from the yc against T plots (see fig. 13 later). Finally, the y values between the equilibrium phases used in the distribution studies are represented. The agreement between the sets of data is excellent. From a comparison of the tension and distribution data for equilibrated systems (fig. 11) we note that the minimum yc occurs at the salt concentration corresponding to the onset of transfer of surfactant to the alkane phase, and hence to the phase inversion of the emulsion formed from these systems. Thermodynamic Treatment of Tension Data A thermodynamic treatment applicable to the systems of present interest has already been de~eloped.~' In particular the equation relating to the variation of yc with salt concentration is where Tc, is the surface excess of the Cl- ion,.fTacl is the mean ionic activity coefficient of the electrolyte and mXa is the total counterion concentration (i.e.salt concentra- tion+c.m.c.). The charge signs on the subscripts have been omitted for clarity. In the derivation of eqn (6) it has been assumed thatf, for surfactant and salt are equal. For a surface containing an adsorbed film of anionic surfactant there is evidence that Tc, is close to zero and slightly negative,28 and for the moment we will ignore the term in Tcl (but see below). Further, we will assume rA is independent of salt concentration. As mentioned, r, does not vary with salt concentration above 0.05 mol dm-3 (corre- sponding to the right-hand side of the curve).For lower concentration r A decreases, but for 0.02 mol dm-3 rA is lowered only by 10-1 5 % . We will report elsewherel5 on the variation of r A with salt concentration and its significance. The curve predicted by eqn (6) (for Tcl = 0) for the present system has been obtained by integrating the equation, which yields for constant rA and T - ycJ = RTT, [In mXa + 2 lnfyacl - +In (c.m.c.)] + B' (7)R. Aveyard, B. P. Binks, S. Clark and J . Meud 137 where B is an integration constant. As shown elsewhere,27 for NaCl eqn (7) may be expressed [using values offf;iaC1 - from ref. (29)] -6.83 x 10-4(lnmN,)3-0.025(lnm,,)2 d In (c.m.c.) d lnm,, ) ln mNa] + B (8) where B is the tension at mNa = 1 mol dmP3.From values of the c.m.c. measured as a function of mNa (see Appendix) we find that d In (c.m.c.)/d In mNa is constant and equal to -0.855. The full curve shown in fig. 10 is obtained from eqn (8) with a value of B = 1.078 mN m-l. Agreement between experiment and eqn (8) is satisfactory, both in terms of the shape of the curve and the salt concentration for a minimum ye. There are, however, significant discrepancies which are presumably associated with the assumption that Tcl is zero, and we now explore this further. Eqn (6) can be expressed asz7 where am and ar, are apparent degrees of dissociation of the surfactant in the micelle and the plane surface, respectively. The quantity ar, is defined as - 2rC,/rA, and a, has been shown by Hal130 to be given by C: ln.fyacl ..=[1+(’ c? Inm,, - )+ d d In lnm,a (c.m.c.)]/[ (2 a lnfyacl)] lnm,, (10) It is seen from eqn (9) that minimum y c is attained when a,, = ar,. When earlier we assumed Tcl = 0 we were setting % = 0 for all mNa.However, we may obtain an expression for o+ by rearrangement of eqn (9), which gives 1 + - - . The ( y c , lnm,,) data are fitted by yJmN m-l = 1.843+ 1.210 Inm,,+0.19905 (In mNa)% (102/r,)/nm2 molecule-l = 41.33 exp ( - 66.13 mNa) + 73.0. (12) (13) and, as reported elsewhere,15 the (rA, In mNa) data can be represented by In fig. 12(a) we show 01, calculated from eqn (10) and % obtained from eqn ( 1 lb(13) as a function of salt concentration. Both a,, and ar, are very small, of comparable magnitude and of opposite sign. Thus, although Tcl is indeed very small [as shown in fig.12(b)] it is not negligible in the context. The interesting observation is that both a, and ar, are zero at mKa = 0.05 mol dm-3, where minimum y c is observed. This is why the curve generated by eqn (8) (for which it is assumed that Tcl and % are zero) gives the minimum at the correct salt concentration. At low salt concentrations a, is positive and the surfactant aggregates are in the aqueous phase. For higher salt concentrations a, is negative, implying that the micelles become positively charged. However, in this range of mNa, the aggregates exist not in the aqueous phase but in the alkane phase, which is a dilute W/O microemulsion, as discussed. The negative values of ar, at low mNa indicate a slight positive adsorption of C1-.The effect of temperature on yc for four salt concentrations is shown in fig. 13. The minimum tension is shifted to higher temperatures as the salt concentration is increased. This is in accord with the earlier observations concerning phase inversion (fig. 1 ) and surfactant transfer between phases (fig. 4). As seen, as the salt concentration is increased,138 0.0: -0.0: Tension Minima in Heptane-Aqueous NaC1-A0 T Systems AOT AOT in water ++ in heptane P 0.02 0.05 2 . 5 2 .o 1.5 1 .o 0 . 5 0 I- 1 0.10 0.02 0.05 In (rnN,/mol dm-3) 0.1 0 Figure 12. (a) Variation of a, (0) and o ~ p (0 j with Naf concentration in the AOT-H,O-heptane system at 25 "C. (bj Tcl (0) and rA (0) as a function of Na' concentration in the same sy s tem. 20 40 60 TI"C Fig. 13. Variation of yc with temperature for four salt concentrations.Curves are for salt concentrations of 0.0342 (@), 0.0512 (O), 0.0856 (a) and 0.1026 (0) mol drnp3.R. Aveyard, B. P. Binks, S. Clark and J. Mead 139 xm is decreased, and when a, = 0 surfactant transfers to the alkane phase. It can be appreciated from fig. 4 that for a given salt concentration, surfactant resides in the oil at low temperature and transfers to the aqueous phase at higher temperatures, and we may suppose that a, becomes more positive with increasing T. Thus, increase in salt concentration and T affect a, in opposite senses, and it can be appreciated why the temperature corresponding to minimum tension increases as the salt concentration is increased. For the presence of a swamping excess of salt, dy,/dT is given byz7 so that T i S , - Si = Sj(Ti N j - Tj") + dy,/dT (15) j in which the Ts are total surface concentrations, S are partial molar entropies in solution, S , is the entropy of micelles containing 1 mol surfactant and Pu is the entropy of unit area of interface; xi denotes summation over all species present and zj summation over all components other than surfactant A, treated as the electrically neutral species.The N j are number of moles of species j in the micelles containing 1 mol A. The first term in parentheses in eqn (14) is the entropy of interface formation (per unit area) AFu, and the second term in parentheses is the molar entropy of micelle formation at the c.m.c., AS,. The significance of eqn (14) and (15) and of dy,/dT can be understood as follows.To form unit area of interface we transfer (Fl +cj T;) mol from the appropriate bulk phases to interface, it being assumed for simplicity that a given component is present in only one of the bulk phases; the entropy change is Si - T i S , -xj rf S j . To form micelles containing T i mol surfactant we transfer (r: +xi rL N j ) mol from bulk, and theentropychangeforthisprocessisT,S, - T i S,-&r% N j S?.Theterm& (Ti N j - A J ) in eqn (1 5) arises from the difference in composition of micelles and interface. Suppose now we remove material from unit area of interface and transfer the TL mol surfactant to micelles. The Tg mol at the interface are associated with r? mol of the other components. In the micelle, however, the ri mol are combined with cj Ti Nj mol of other components, so that Ej (ri N j - rj") mol are removed from bulk with an accom- panying entropy loss of Sj(ri N j - rf).The overall entropy change for the transfer is therefore r A S , - Si - xj Sj(Ti N j - r;), which as seen from eqn (1 5 ) is simply dy,/dT. At the temperature r* corresponding to minimum yc this entropy change is zero as a result of the equality (16) rA S , - s: = c sj(r% N~ - rj"). dyc/d T = rh ASm - AS: (17) so that at T = r* r A AS, = AS:. (18) -ASm z RT[d In (c.m.c.)/dT]. (19) j We note also that eqn (14) may be written The micellisation entropy is given to a good approximation by3" Values of rAAS, (obtained using the temperature varation of the c.m.c. given in the Appendix for AOT in 0.085 mol dmP3 NaCl together with T i = 2.3 x lop6 mol mP2, which is equivalent to 73 A2 molecule-') are shown in table 2 for T > P, where surfactant is present only in the aqueous phase.dy,/dTis only 2 or 3% of the magnitude of TL ASn, and AS:, which are thus almost equal. It follows that the tension minimum with respect to temperature arises from a very fine balance between the entropies of surface and micelle formation.140 Tension Minima in Heptane-Aqueous NaC1-AOT Systems Table 2. Entropies of micelle and surface formation in the system AOT-heptane-0.085 mol dm-3 NaCP ~ T/"C lo6 dy,/dT - lo4 ri ASm - 10' APu 50 2.2 1.30 1.32 52 3.8 I .42 1.46 54 4.6 1.53 1.58 56 4.4 1.65 1.69 a All units are J m-' K-l. -9 - 8 -7 -6 In ( [ AOTj/iiiol dni-') 7 h m I r - 3 - 2 -? 0 - "1 r r -I b e - - I I -2 9 - 3 In (mN,/mol dm-3) Fig.14. (a) Tensiometric determination of the c.m.c. of AOT in the presence of NaCl and excess heptane phase at 25 "C. Points 0, a, a, and 0 are for salt concentrations of, respectively, 0.0257, 0.0513, 0.0856 and 0.1369 mol dm-". (b) Dependence of c.m.c. on total counterion concentration at 25 "C. Points (0) are from tensiometric measurements as in (a); points (a) are equilibrium aqueous-phase concentrations for salt concentrations where aggregation occurs in heptane.R. Avejmrd, B. P. Binks, S. Clark and J . Mead 141 The authors express their gratitude to the British Petroleum Company (B.P.) for an Extramural Research Award and for the award of a B.P. Research Studentship to B.P.B. Appendix Critical Micelle Concentrations We have determined critical micelle concentrations of AOT in heptane-aqueous NaCl systems tensiometrically both as a function of salt concentration and temperature.The c.m.c. is taken as the aqueous surfactant concentration at which y just attains the constant value yc. As explained earlier, aggregation may occur in either the oil or the aqueous phase at this point. Some tension data are shown in fig. 14(a) and the variation of c.m.c with salt concentration is depicted in fig. 14(b). The change in c.m.c. (in 0.085 mol dmP3 NaC1) with T (not shown) was obtained from four values between 298 and 333 K; the data are fitted by In (c.m.c.) = 62.3889+0.5947T-2.1803 x 10-3T2+2.6671 x IOp6T3. (20) In the distribution studies we determined equilibrium concentrations of AOT in aqueous solution (in contact with heptane) as a function of salt concentration.It was found that, for high salt concentrations, the AOT concentrations were ca. 10 % higher than the c.m.c. values determined tensiometrically, and yet no micelles were detectable by P.C.S. As already discussed, the y values determined in the spinning-drop experiments appear to be equilibrium values (figure lo), and at present we have no explanation for this discrepancy. A similar situation has been observed by Cazabat et al.31 in systems containing sodium dodecyl sulphate. However, the logarithm of the equilibrium aqueous-phase concentration of AOT is linearly related to lnm,, [fig. 14(b)], the slope of the line being equal within experimental error to d In (c.m.c.)/d In m,,, so the analysis using eqn (6) and (8) is unaffected.References 1 J. T. G. Overbeek, P. L. De Bruyn and F. Verhoeckx. in Surfuctunts, ed. Th. F. Tadros (Academic Press, London, 1984), chap. 5. 2 H. Kunieda and K. Shinoda, Bull. Chem. SOC. Jpn, 1982, 55, 1777. 3 K. S. Chan and D. 0. Shah, J. Dispersion Sci. Technol., 1980, 1. 55. 4 E. 1. Franses, J. E. Puig, Y. Talmon, W. G. Miller, L. E. Scriven and H. T. Davis, J. Phy.7. Chem., 1980, 5 E. I. Franses, Y. Talmon, L. E. Scriven, H. T. Davis and W. G. Miller, J . Colloid Interfuce Sci., 1982, 6 J. E. Puig, E. 1. Franses and W. G. Miller, J. Colloid Interface Sci., 1982, 89, 441. 7 M. Dupeyrat, L. Minssieux and A. El. Naggar, European Symp. Enhanced Oil Rerouery (Edinburgh, 8 E. I. Franses and T. J. Hart, J. Colloid Interface Sci., 1983, 94, 2. 9 R. Aveyard and D. A. Haydon, Trans. Faradaq, Soc., 1965, 61, 2255. 84, 1547. 86, 449. 1978), p. 161. 10 R. Aveyard and S. M. Saleem, J. Chem. SOC., Furuduy Trans. I , 1976, 72, 1609. 11 V. W. Reid, G. F. Longman and E. Heinerth, Tenside, 1967, 4, 292. 12 K . Shinoda and K. Ito, J. Phys. Chem., 1961, 65, 1499. 13 K. Fischer, Angew. Chem., 1935, 48, 394. 14 D. E. Koppel, J. Chem. Phys., 1972, 57, 4814. 15 R. Aveyard, B. P. Binks and J. Mead, to be published. 16 K. Shinoda, M. Hanrin, H. Kunieda and H. Saito, Colloids Surf, 1981, 2, 301. 17 K. Shinoda and H. Kunieda, J. Colloid Interface Sci., 1973, 42, 381. 18 H. Saito and K. Shinoda, J. Colloid Interface Sci., 1970, 32, 647. 19 C. Tanford, Physical Chemistry of Macromolecules (Wiley, New York, 1961). 20 J. W. Mehl, J . L. Oncley and R. Simha, Science, 1940, 92, 132. 21 P. C. Hiemenz, Principles of Colloid and Surface Chemistry (Dekker, New York, 1977). 22 S. Mukherjee, C . A. Miller and T. Fort, J. Colloid Interface Sci., 1983, 91, 223.142 Tension Minima in Heptane-Aqueous NaCI- .4 0 T Sj-stems 23 J. D. Nicholson and J. H. Clarke, Surfactants in Solution, ed. K. Mittal (Plenum Press, New York, 24 B. H. Robinson, C . Toprakcioglu, J. C . Dore and P. Chieux, J. Chem. Soc., Furaday Trans. I , 1984, 25 M. Zulauf and H-F. Eicke, J . Phys. Chem., 1979, 83,40. 26 E. Matijevic and B. A. Pethica, Trans. Faraday Soc., 1958, 54, 1382. 27 R. Aveyard, B. P. Binks and J. Mead, J. Chem. Soc., Faraday Trans. 1, 1985. 81, 2169. 28 K. Tajima, Bull. Chem. Soc. Jpn, 1971, 44, 1767. 29 R. A. Robinson and R. H. Stokes, Electrolyte Solutions (Butterworths, London, 1959). 30 D. G. Hall, in Aggregation Processes in Solution, ed. E. Wyn-Jones and J. Gormally (Elsevier, Amster- 31 A. M. Cazabat, D. Langevin, J. Meunier and A. Pouchelon, Ad[,. Colloid Interface Sci., 1982, 16, 175. 1984). 80, 13. dam), chap. 2. Paper 51552; Received 1st April, 1985

ISSN:0300-9599    

DOI:10.1039/F19868200125

出版商:RSC    

年代:1986

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I , 1986, 82, 143-156 Protonic and Ionic Conduction in Lysozyme Hydration and Field-dependent Effects Hywel Morgan and Ronald Pethig* School of Electronic Engineering Science, University College of North Wales, Bangor, Gwynedd LL57 I UT Solid-state conduction and dielectric measurements are reported for lysozyme as a function of hydration and applied electric field. Mass spectrometry and electron-probe microanalysis have been used to detect protonic and ionic conduction processes. For electric fields < ca. 8 x lo5 V m-l the electrical conductivity in lysozyme is predominantly protonic for all hydrations above 5 wt %. For higher electric fields Poole-Frenkel-dominated ionic transport processes become progressively more dominant as the hydration content is increased.The observation of a pronounced isotope (deuteration: effect and the result of prolonged field cleaning of the test samples support these conclusions. The low-frequency dielectric dispersion (known as the n- dispersion) commonly observed for protein samples has been found to be associated with interactions between ions and the metal electrodes. ~ The solid-state electrical conduction properties of proteins and related materials has been the subject of many experimental and theoretical studies. The main stimulus for this was Szent-Gyorgyi's proposal1 that delocalized electronic states in proteins may have a biological role. The hydration-dependent semiconductive (Arrhenius-type) properties of protein powders were first established in the laboratories of Eley23 and Rosenberg4 more than 20 years ago.However, as indicated by recent review^,^^ there is as yet no clear understanding of either the processes by which protein structures sustain electrical charge transport or of the nature of the dominant charge carriers. However, some aspects do appear to be clear, e.g. the fact that proteins isolated in their pure and dry condition exhibit a negligibly small intrinsic semiconductivity.23 :3 The role of adsorbed water is therefore of fundamental importance. Also, the~retical~-~ and microwave Hall-effect studies1*$ l1 indicate that electron transport in protein structures is more likely to occur through hopping or phonon-assisted tunnelling mechanisms than as a delocalized wave-like process. The main reason for this is that the broad valence and conduction bands of extended electronic states expected from earlier 12-14 theoretical considerations of periodic homopolypeptide structures become disrupted for the aperiodic structures of natural l6 Electrical studies17 on single crystals of lysozyme have indicated that protonic transport may be the most dominant conduction process, and measurements1s-20 on natural and deuterated lysozyme powders have shown that both ionic and protonic conduction mechanisms contribute to the overall dielectric processes. The measurements reported here have been made on freeze-dried powders of lysozyme, and this has made it possible to investigate the dielectric and electrical conduction properties as a function of hydration of the powders.Lysozyme is a particularly stable enzyme.21 There is evidence22 to show that its conformation in a dry powder preparation is similar to that when in solution, and when partially dehydrated it continues to exhibit enzymatic activity.23 A wide range of physical measurements have been applied to lysozyme, and the n.m.r., e.s.r., infrared, Raman, heat-capacity and dielectric properties as a function of hydration have recently been reviewed.24 Our present studies provide new information concerning the influence of high electric fields and hydration on the 143144 Conduction in Lysozyme a i r inlet - remote r.f. sa It solution Anavac 2 head mass spectrometer -& sample I L high - v a c p u m p s Fig. 1. Schematic outline of the apparatus used for the measurement of protonic conductivity. transport of ions and protons in lysozyme in particular, and for protein structures in general.Apart from the relevance of this to the general area of biophysics, such information is also of relevance to the development of enzyme electrodes and biosen~ors.~~ In this respect, the fact, for lysozyme at least, that full solution conditions are not necessary for enzymatic activity is of particular importance. Experiment a1 Egg-white lysozyme (grade 1) was obtained from Sigma Chemicals and compressed at 1.3 x lo8 Pa into discs of surface area 1.27 x m2. For some of the studies the samples were first deuterated by dissolvingin deuterium oxide (99.8 atom :< D, Sigma Chemicals) for 48 h before freeze drying. The disc samples were retained between two spring-loaded polished copper electrodes in a temperature-controlled and vacuum-tight electrical measurement cell, which was incorporated into a vacuum and mass-spectrometer system as shown in fig.1. Voltages ranging from 200 mV to 2 kV were applied to samples of thicknesses ranging from 0.5 to 3 x lop3 m, and sample currents were measured using a Keithley electrometer (model 602 or 610) and a chart recorder. Any deuterium gas emitted from the deuterated samples at the cathodic electrode was absorbed by the palladium foil, of weight 1 g and surface area 6.5 x m2. The amount of deuterium absorbed was later determined by isolating the palladium foil from the electrical conduction cell, evacuating to lop4 Pa, heating to 1143 K and measuring the deuterium released over 5 h using an Anavac 2 (VG Ltd) mass spectrometer.The system had previously been calibrated from electrolysis measurements on D20. The partial pressure of either H 2 0 or D20 in the conduction cell was controlled using calibrated H,O or D20 saturated-salt solutions and a temperature bath. H 2 0 and D20 sorption isotherms were obtained using a Cahn electrobalance in a temperature-controlled vacuum system. A Jeol JXA-3A electron-probe X-ray microanalyser was used to inspect the sample surfaces for any products produced as a result of solid-state electrolysis. Dielectric measurements of the complex permittivity were made in the frequency range Hz to 100 kHz, using voltage-step response and a.c. bridge techniques described elsew here. 26H.Morgan and R. Pethig 145 -5 n d 5 - -7 --. OD -9 16 13.1 10.2 I I 2 3 log ( VIV, Fig. 2. Current-voltage characteristics typically observed for compressed lysozyme samples at different hydrations (values in wt ”/, H,O), showing a well defined transition voltage V, separating the low- and high-field conduction regions. The sample thickness for these results was 9.7 x m. Results and Discussion For constant hydration and temperature, the current-voltage characteristics were typically as shown in fig. 2, with near-ohmic behaviour at low voltages changing to a Vn law at a well defined voltage &, with n having a value of 2.2kO.l. Values for the exponent n of this order are often indicative of space-charge-limited conduction, and the current density Jshould vary2’ with voltage Vand sample thickness d as V2dP3.However, our measurements for different samples of thicknesses ranging from 3 x 1 0-4 to 2 x 1 OP3 m have given a thickness dependence of V2dP2, indicating the space-charge-limited con- dition is not being observed for the results of fig. 2. V, varies linearly with the reciprocal of the 100 kHz relative permittivity value, as shown in fig. 3 . Also, as shown in fig. 4, at high fields the conductivity and steady-state current vary as the square root of the field. This behaviour is consistent with either the Schottky or Poole-Frenkel According to the Schottky law, the current density J varies with the applied field E as J = AT2exp - ( w - ~ E + ) / w (1) where A is the emission coefficient and Wis the zero-field barrier at the electrode-sample interface.The field-lowering parameter is given by p = (q3/4nE0E,)+ with q being the electronic charge, E , the high-frequency relative permittivity of the sample and E, the permittivity of a vacuum.27 The bulk analogue of the Schottky effect is the146 Conduction in Lysozyme 9 00 700 vt 5 00 300 0 0 0.1 0.2 1 kr Fig. 3. Relationship between the transition voltage V, and the sample relative permittivity E, measured at 100 kHz. -9.5 n I - E ?. -9.0 v on - -10.2 I I 1 I 1 I I I I 1 -6.0 5 2! 9 W h v -7.0 -8.2 12 16 20 24 28 30 fieldi/Vt rn-; Fig. 4. Schottky ([7) and Poole-Frenkel (e) plots [eqn ( 1 ) and (2), respectively] for lysozyme samples at 294 K: (a) 6.6 wt % H,O, (h) 7.8 wt % H,O.H . Morgan and R. Pethig 147 Poole-Frenkel effect of field-assisted thermal activation of charge carriers over coulombic barriers.For the case 29 where these coulombic barriers are associated with trapping sites of depth W, the conductivity varies with the local field E , according to the relationship (2) A clear-cut distinction between these two processes is often not possible, but some indication can be obtained by determining the values for E , derived from the slopes of the linear high-field regions. Values derived for E , from the Schottky plots of fig. 4 are 0.7 and 1.0 for hydrations of 6.6 and 7.8 wt % , respectively. These permittivity values are not physically realistic. The corresponding relative permittivity values derived from the Poole-Frenkel plots are 4.0 and 6.8, respectively, which are in reasonable agreement with the values obtained at 100 kHz at these hydrations.Also, it has been observed in this laboratory30 and elsewhere31 that the steady-state conduction and dielectric properties of protein samples are not strongly dependent on the metal (copper, gold, silver and aluminium) used for the electrodes. (If hydrogen-saturated palladium-black electrodes are used, however, then as for cyclodextrin samples, larger conductivity values are obtained, consistent with the effect of proton injection from the anodic electrode.) These results indicate that the effects shown in fig. 2 and 4 are associated with bulk conduction processes. The sharp transitions in the current-voltage characteristics shown in fig. 2 suggest that two separate conduction mechanisms operate at the low- and high-field regions.with each mechanism involving a different type of charge carrier. Teda et aJ.33 have shown for a single charge carrier that there is a gradual change in the current-voltage characteristic as ohmic conduction becomes modified by the Poole-Frenkel effect. As will be described later, there is good evidence that the low-field conductivity is protonic and that Poole-Frenkel-dominated ionic transport processes begin to occur at the higher field strengths. The total conductivity can thus be represented by the sum of two components ol7 = a,+o, where op is the protonic conductivity, and the ionic conductivity is given by The permittivity factor &h is included in this equation to allow for the effect4 of hydration in lowering the effective value of the potential-energy barrier from its dry-state value of Wo.At the transition voltage &, corresponding to a macroscopic field Et = &/d where d is the sample thickness, the conductivities op and oi are of equal magnitude, and so o = oo exp -( W - / I E ~ ) / L T . oi = 0, exp ( - W,/e, k T ) exp ( /IEi/k T). (3) from eqn (3) where the identities EL = aEt and Eh = Y E , have been used to relate the local values of the electric field and relative permittivity to their macroscopic values. From the measured values of op and the extrapolated value for 0, at zero field, at room temperature it may W,/ye, B kT In o,/o,. be concluded that Thus to a good approximation the transition voltage is given by This gives a reciprocal relationship between I/t and the sample permittivity E,, in agreement with the effect shown in fig. 3.The straight-line slope of fig. 3 has a value of 1500 V, and the sample thickness was 6.7 x lo-* m. From eqn (4) this gives W, = 0.1 1 (a?)& eV.148 Conduction in Lysozyme - 4 - 6 -8 h d I - 3 m -. v 3L -10 - -12 -14 0 40 80 120 160 H O/ 1 y soz y in e Fig. 5. Hydration-dependent conductivity for a lysozyme sample measured at low and high fields (@. 290 and 0. 1.2 x lo6 V m-l). (40 sorbed H,O molecules per lysozyme corresponds to a hydration of 4.9 wt .) The simple result of Lorentz gives a = (~,+2)/3, and since the mobile ions are likely to he located near the hydrophilic groups of the protein structure (giving a local polarisability larger than the average bulk value) then the value for y is likely also to be greater than unity.Taking into account the simplifying assumptions in the Poole- Frenkel model used here, namely a two-dimensional rather than three-dimensional analysis and the assumption of a simple coulombic barrier, this result for W, can be considered to be of the correct order of magnitude and is comparable to the value of 0.55 eV for m were examined at the same time under identical hydrations at 294 K. To one sample a voltage of 0.28 V was applied and to the other 1200 V, corresponding to the ohmic and Poole-Frenkel conduction regions: respectively. The conductivity results obtained are shown in fig. 5 , and the significant field-dependent difference in conductivity at the higher hydrations does not appear to have been described previously in the literature.The number of adsorbed water molecules given in fig. 5 was derived from a hydration isotherm of the type shipwn in fig. 6. Before the electrical measurements were made the sample surfaces appeared white, but ‘Two samples of compressed lysozyme of thickness 1 xH. Morgan and R. Pethig 149 160 120 - 2 - 0 E . x 0, 8 0 3: a -e 2 0 C 40 c13 0 0 0.2 0.4 0.6 0.8 H20/D20 partial pressure Fig. 6. H20 (a) and D20 (0) sorption isotherms for field-cleaned lysozyme at 294 K . Values derived for the B.E.T. parameters*" h, and C are: H 2 0 isotherm (h, = 5.9, C: = 20.7) and D,Q isotherm (h, = 6.1, C = 29.0). after obtaining high-field results such as those of fig. 5 the surface previously in contact with an anodic copper electrode commonly assumed a blue-green colouration.Electron- probe microanalysis of this surface has shown that this surface discolouration is associated with the presence of sulphur and chlorine copper salts, indicating that sulphate and chlorine ions had contributed to the conduction current. The cathodic sample surface exhibited no visible discolouration, and apart from copper no other metallic ions were detected. Copper sulphate appeared to be the dominant impurity on the anodic surface, and the amount of this was determined by dissolving the top layer of the surface in a known volume of deionised water, centrifuging and comparing the visible absorption spectrum against calibrated copper sulphate solutions. The amount of sulphate ions so detected was found to account for ca.80% of the total charge passed through the sample for the conductivity measurements at the highest hydration levels. The effect of migrating ions also appears in the high-field current-time characteristics of hydrated samples, as shown in fig. 7 where a current peak appears ca. 24 h after the application of the field. The location in time of this current peak was found to be linearly dependent on voltage, but almost independent of hydration. This peak has similarities to the transient space-charge-limited current (s.c.1.c.) peaks which have been observed for p~lyethylene~~ and hydrated poly(~inylchloride).~~ According to the s.c.1.c. analysis, where there is continuous injection of carriers into the sample, the mobility of the charge carriers can be derived from the expression p = 0.787 d2/Vtp where Vis the applied voltage and t , is the peak time.For the sample of fig. 7 this gives a mobility value of 1.8 x m2 V-l s-l, which is of the same order as the values derived for p~lyethylene~~ and poly(~inylchloride).~~ In the absence of space-charge effects and150 Conduction in Lysozyme 0 0 0 0 - 6 t I I I I I 1 0 1 .o 1.5 Fig. 7. Double-log plot of the current-time characteristic for a lysozyme sample subjected to a field log ( t l h ) of 3 x lo6 V rn-l at 12 wt hydration. diffusion the mobility is given by d 2 / Vt,, and the peak time t, corresponds to the transit time of the injected carrier front. However, we consider that such an analysis is not pertinent to our experiments. Instead, 38 associated with the thermal detrapping of bulk charges may be more relevant, since we appear to be dealing with the diffusion of impurity ions originating in the sample bulk and not s.c.1.c.effects. If samples that have been subjected to high fields and high hydrations for long periods of time (> 20 h) have the discoloured surface layer (described above) removed, then it is found l9 on subsequent conduction measurements that the sample conductivity has been reduced by orders of magnitude. This result is therefore equivalent to a field cleaning process in which ions in the bulk of the sample are removed, and the effect of this on the conductivity is shown in fig. 8 for three deuterated samples of similar thickness with applied voltages of 400 V. Sample (a), exhibiting the largest conductivity in fig.8, had received no field cleaning treatment, whereas sample (6) had received 240 h of field cleaning (equivalent to the passage of 17 C through the sample), and for sample (c) the corresponding field cleaning time was 500 h. Our previous worklg has shown that Sephadex gel filtration also results in a lowering of sample conductivity, but the extent of this is much less. Also included in fig. 8 are the results obtained using the mass spectrometer to detect the emission of deuterium gas from the sample-cathode interface. The deuterium ion conductivity for constant levels of deuteration was determined by applying a constant voltage to the sample and monitoring the average steady-state current over a period of 18 h. The amount of deuterium gas emitted and absorbed by the palladium foil during this period was later measured by the mass spectrometer.From this the average deuterium ion current that had passed through the sample could be determined. The effective deuterium ion, hence also the protonic, conductivity was found to be independent of any field cleaning treatment, and the results shown in fig. 8 indicate that for very pure samples the conductivity can be attributed almost entirely to the transport of protons. For conductivities below 3 x S m-l the sensitivity of measurement was such that entirely reproducible values for the protonic conductivity component could not be obtained for current time lengths of 18 h. Measurements were also made at high hydration levels for an applied voltage of 0.28 V (i.e.less than the dissociation voltage of water at normal temperatures), and the observed conductivity was found to be solelyH. Morgan and R. Pethig 151 - 8 -9 -10 h -11 E - I Vl 1 v M 2 -12 - 1 3 - 1 4 -15 -1 1 I I 1 1 0 20 LO 60 80 100 120 140 no. of sorbed D,O molecules Fig. 8. Variation of the high-field conductivity of deuterated lysozyme as a function of D,O content for: (a) no field cleaning, (b) 240 h field cleaning and (c) 500 h field cleaning. 0, Deuterium ion conductivity determined using the apparatus shown in fig. 1. (40 sorbed D,O molecules per deuterated lysozyme corresponds to 5.0 wt % D,O content.) attributable to protonic conduction. This last result confirms the conclusion regarding the protonic nature of the low-field conductivity. More significantly, it indicates that the emitted deuterium gas does not primarily arise from the products of the electrolysis of D20 and that the mobile protons originate from the hydrated protein structure.The H 2 0 and D20 isotherms obtained at 294 0.5 K for lysozyme samples are shown in fig. 6. The D,O isotherm was obtained on deuterated lysozyme that had been field-cleaned in a D20 partial pressure of 0.85. In this way the isotherm should not include weight increases associated with labile hydrogens in the protein structure having been replaced with deuterium atoms. Our previous workl8 has shown that ca. 2.1 hydrogen atoms per amino-acid residue in the lysozyme structure are exchanged. From fig. 6 it can be seen that D20 appears to bind more strongly to lysozyme than does H,O, suggesting that a deuterium bond to the protein structure is stronger than a hydrogen bond.Dyke et aZ.39 in their studies of the radiofrequency and microwave spectra of the152 Conduction in Lysozyme -1 0 -12 - - I E m D w . v - -1 4 -16 0 4 0 80 120 no. of sorbed D,O/H,O molecules Fig. 9.294 K conductivity of normal (0) and deuterated (a) field-cleaned lysozyme as a function of H,O and D,O content, respectively. K = 0 states of (HF), and (DF), and their gaseous mixture, could find evidence for the existence of the HFDF isomer but not for DFHF. This implies that the deuterium bond is energetically more favourable than the hydrogen bond. The reason for this will be associated with the deuterium bond having the lower zero-point energy because of the larger mass of the deuterium atom.The sorption isotherms of fig. 6 also reflect the difference between the two competing processes: (H,O-HOH) + protein + H,O + (HOH-protein) (D,O-DOD) +protein + D,O + (DOD-protein). and The effect of deuterium substitution on the interaction between water molecules as well as on the HOH-protein adduct is therefore involved. Values obtained for the monolayer parameter h , and the monolayer binding affinity C of Brunauer et are given in the legend of fig. 6. The results obtained for the steady-state conductivities of normal and deuterated field-cleaned lysozyme, as a function of H,O and D,O contents, respectively, are shown in fig. 9. A significant isotope effect is apparent, particularly at the lower hydration regions.This result, which has been found to be consistently reproducible, is in agreementH. Morgan and R. Pethig 153 24 - 16 - f" 8 - a - 4 - 3 - 2 log (frequency/Hz) Fig. 10. Frequency dependence of the imaginary component E" of the complex permittivity for lysozyme with 11 wt "/d water content: ( a ) no field cleaning and (b) 500 h field cleaning. with our previous workls and adds further support to the conclusion that for field-cleaned samples the dominant conduction process in lysozyme involves the transport of protons. Deuterium-exchange rneasurement~~~ for insulin have indicated that the most labile hydrogens are those (OH and NH) associated with side chains of the basic and acidic residues, together with accessible backbone imide hydrogens not involved in the hydrogen-bond stabilising of a-helical or P-sheet conformations.If this can be taken as a guide to the behaviour of lysozyme, then from its amino-acid composition and conformationz1 an estimated maximum of around 265 labile hydrogens may be deduced, which compares well with our earlier experimental result.ls There is good evidence24 that the first water molecules to bind to lysozyme interact with the ionisable carboxylic and basic groups, and that proton redistribution occurs at a hydration level as low as 5 wt % . Since the pK values of the basic and acidic groups in a protein are of the order 0.4 to 0.6 pH units higher in DzO than in H20,42 then the number of protons released from the acidic groups should be lower in D,O vapour. This reduction in the number of available mobile protons, together with the lower mobility of heavier deuterium ions, could account for the isotope effect shown in fig.9. Effects associated with differences in protein conformation may also be occurring. Dielectric measurements in the frequency range lop4 Hz to 1 kHz for hydrated proteins have revealed the existence of two dielectric relaxations, designated as the R- and a-dispersions. The a-dispersion is considered43- 44 to be associated with a bulk effect closely related to the steady-state conduction processes, and we will not consider this here in any more detail. The Q-dispersion, which occurs at a frequency lower than the a-dispersion, has been considered43 to be associated with polarisations occurring at the electrode-sample interfaces, and in particular1g with the build-up of impurity ions at the electrodes.Recent t h e o r i e ~ ~ ~ , ~ ~ have suggested that there is a common basis for the R- and a-dispersion processes and that together they form a so-called anomalous dispersion arising from non-conductive long-range charge transport. A difficulty we have with such a viewpoint is that we find when using blocking electrodes, such as mica or PTFE, that the R-dispersion disappears whereas the a-dispersion remains unchanged. A typical plot of the imaginary component E" of the complex permittivity for untreated lysozyme is shown in fig. lO(a), in which regions of both the R- and a-dispersions can be seen. A typical result obtained for a lysozyme sample that has been extensively field-cleaned is given in fig.10(b) and shows that the R-dispersion has disappeared as a result of this treatment. The a-dispersion can be seen to be reduced in magnitude and to occur at a lower frequency, in accordance with earlier observations4*3 47 of the relationship between conductivity and the dielectric relaxation process. This disappearance of the R-dispersion154 Conduction in Lysozyme as a result of the removal of impurity ions by field cleaning strongly suggests that its existence is associated with the drift of such ions towards, and their resulting interaction with, the metal electrodes. Several charge-transfer processes could give rise to the R-dispersion. For example, a non-Faradaic process of the form Cu- repels SO!- Cu+ attracts SO:- could occur at the electrode-sample interfaces and would be equivalent to the discharging and charging of electrical double layers.We have already described how an electrochemical process of the form CuTT + SO:- + CuSO, + 2e-- appears to be significant, and several other electrochemical interactions involving impurity ions are also possible. Conclusions The solid-state electrolysis measurements reported here are similar in principle to those described by Powell and Ro~enberg~~ for various biomacromolecules (but not including lysozyme). For hydrations in the range 6-50 wt "/o , all the samples tested by these authors appeared to be mixed conductors in that both electronic and protonic charge carriers were concluded to contribute to the total conductivity. From our own measurements we are not able to draw the same conclusions for lysozyme. Under conditions of high electrical field strength and hydrations, we find that the electrical properties of our lysozyme samples are dominated by impurity ions.As these ions are removed from the sample by gel filtration or field cleaning, protonic conduction becomes more dominant (fig. 8) for hydrations above 5 wt "/o . Apart from the measurements described here, we have also investigated the electrical properties of dry lysozyme in high vacuum conditions. For fields below lo6 V m-l no hydrogen gas was detected, so the small conductivity (ca. S m-l) could conceivably be electronic in nature. For larger fields, hydrogen and other gasses such as NH,, CO, CO, and 0, for example, were detected by the mass spectrometer.This most likely resulted from the oxidation of the protein structure owing to the high electrical stress, involving such reactions as Cn€12nOn + nH,O + nCO, + 4nH+ + 4ne- for example, and cannot be taken as evidence for protonic or ionic conduction. (Even when 'dry' there are ca. four or five water molecules strongly bound within the structure of each lysozyme molecule.) By comparing the results of fig. 5 and 8 it may also be concluded that for low fields (i.e. less than ca. 8 x lo5 V m-l) the conductivity is mainly protonic at all hydration levels above 5 wt % , even if the samples contain impurity ions. As the field is increased, we picture the potential-energy barriers that limit the conduction of the loosely trapped impurity ions being reduced by the Poole-Frenkel effect.This results, at the higher hydrations, in the ionic conductivity increasing to a level that exceeds the protonic conductivity. The observation of the isotope effect in fig. 9 and the recent observation by Careri et alS2O that the a-dispersion in hydrated lysozyme is pH-dependent and is altered by deuteration, also provides strong evidence for the existence of protonic conductivity. Careri et a1.20 conclude that the proton flow is cooperative, and that this may be of particular significance for the enzymatic functioning and other properties of proteins. In earlier on bovine serum albumin and lysozyme we suggested that the protonic conduction involves the percolation of protons along networks of protein-water and water-water hydrogen bonds. The protons were considered to originate from the ionizable acid side groups, and their mobility was suggested also to be influenced by the ease with which sorbed water molecules could diffuse about the protein surface so as to achieve the optimum configurations for proton transfer.Depending on the appliedH. Morgan and R. Pethig 155 field strength, the limiting protonic conductivity at high hydrations is of the order 10-ll-lO-lo S m-l. If, as a first estimate, we assume a protonic mobility p of the order lop7 m2 V-l s-l (as for ice)5* then using the relationship ~7 = Nqp, such a conductivity requires a proton concentration N of ca. 6 x 1015 mp3. This corresponds to one mobile proton per lo1* lysozyme molecules in the sample and, even allowing for a smaller mobility, this indicates that more than a sufficient number of protons are available for the conduction mechanisms described in this work.This work is supported by the National Foundation for Cancer Research (USA). We thank Mr George T. Stevens for constructing the glassware and assisting in its design, and Prof. G. Careri for sight of ref. (20) before its publication. References 1 A. Szent-Gyorgyi, Nature (London), 1941, 148, 157. 2 D. D. Eley, Horizons in Biochemistry, ed. M. Kasha and B. Pullman (Academic, New York. 3 D. D. Eley, Organic Semiconducting Polymers, ed. J. E. Katon (Marcel Dekker, New York, 4 B. Rosenberg, J . Chem. Phys., 1962, 36, 816. p. 341. p. 259. 5 R. Pethig, Dielectric and Electronic Properties of’BioIogica1 Materials (J.Wiley, Chichester. 1979). 6 R. Pethig, Noncrystulline Semiconductors, ed. M. Pollak (CRC Press, Boca Raton), in press. 7 M. Kertesz, J. Koller and A. Azman, Phys. Rev. B, 1978, 18, 5649. 8 E. G. Petrov, Int. J. Quantum Chem., 1979, 16, 133. 9 S. Larsson, Int. J. Quantum Chem., Quantum Biol. Symp., 1982, 9, 385. 10 D. D. Eley and R. Pethig, Discuss. Faraduy Soc., 1971, 51, 164. 11 T. E. Cross and R. Pethig, Int. J. Quantum Chem., Quantum Biol. Symp., 1980, 7, 385. 12 M. G. Evans and J. Gergely, Biochim. Biophys. Acta, 1949, 3, 188. 13 M. Suard, G. Berthier and B. Pullman, Biochim. Biophys. Acta, 1961, 52, 254. 14 J. Ladik, Nature (London), 1964, 202, 1208. 15 S. Suhai, J. Kaspar and J. Ladik, Int. J. Quantum Chem., 1980. 17, 995. 16 R. S. Day, S. Suhai and J. Ladik, Chem.Phys., 1981, 62, 165. 17 M. Ataka and S. Tanaka, Biopolymers, 1980, 19, 669. 18 J. Behi, S. Bone, H. Morgan and R. Pethig, Int. J. Quantum Chem., Quantum Biol. Symp.. 1982,9, 367. 19 H. Morgan and R. Pethig, Int. J. Quantum Chem., Quantum Biol. Symp., 1984, 11, 209. 20 G. Careri, M. Geraci, A. Giansanti and J. A. Rupley, Proc. Natl Acad. Sci. USA, 1985, 82, in press. 21 T. Imoto, L. N. Johnson, A. C. T. North, D. C. Phillips and J. A. Rupley. in The Enzymes, ed. P. D. Boyer (Academic Press, New York, 1972), p. 665. 22 G. Careri, A. Giansanti and E. Gratton, Biopolymers, 1979, 18, 1187. 23 P-H. Yang and J. A. Rupley, Biochemistry, 1979, 18, 2654. 24 J. L. Finney and P. L. Poole, Comments Mol. Cellular Biophys., 1984, 2. 129. 25 C. R. Lowe, Biosensors, 1985, 1, 3. 26 P. Carnochan and R. Pethig, J. Chem. Soc., Faraduy Trans. I , 1976, 72, 2355. 27 J. G. Simmons, J. Phys. D, 1971, 4, 613. 28 J. Frenkel, Phys. Rec., 1938, 54, 647. 29 C. A. Mead, Phys. Rev., 1962, 128, 2088. 30 H. Morgan, Ph.D. Thesis (University of Wales, 1985). 31 D. D. Eley and P. W. Thomas, Trans. Furaday Soc., 1968, 64, 2459. 32 S. Bone and R. Pethig, Int. J. Quantum Chem., Quantum Biol. Symp., 1983, 10, 133. 33 M. Ieda, G. Sawa and K. Sousuke, J. Appl. Phys., 1971. 42, 3737. 34 G. H. Bardelmeyer, Biopolymers, 1973, 12, 2289. 35 T. Mizutani and M. Ieda, J. Phys. D, 1979, 12, 291. 36 M. Onoda. H. Nakayama and K. Amakawa, Jpn. J . Appl. Phys., 1981, 20, 861. 37 A. von Hippel, E. P. Gross, J. G. Jelatis and M. Geller, Phys. Rev., 1953, 91. 568. 38 R. B. Comizzoli and F. K. Manasse, J. Appl. Phys., 1968, 39, 4868. 39 T. R. Dyke, B. J. Howard and W. Klemperer, J. Chem. Phys., 1972, 56, 2442. 40 S. Brunauer, P. H. Emmett and E. Teller, J . Am. Chem. Soc., 1938, 60, 309. 41 K. Linderstrom-Lang, in Symposium on Peptide Chemistry (Taylor and Francis, London, 1955), p. 1. 42 J . P. Klinman, Adti. Enzymol., 1978, 46, 41 5. 43 D. D. Eley, N. C. Lockhart and C. N. kchardson, J. Chem. Soc., Faraday Trans. 1. 1979, 75, 323. 44 J. Eden, P. R. C. Gascoyne and R. Pethig, J . Chem. Soc., Faraday Trans. I , 1980, 76, 426.156 Conduction in Lysozyrnp 45 L. A. Dissado and R. M. Hill, J. Chem. SOC., Faraday Trans. 2. 1984, 80, 291. 46 M. Shablakh, L. A. Dissado and R. M. Hill, J. Bid. Phys., 1984, 12, 63. 47 P. R. C. Gascoyne and R. Pethig, J. Chem. SOC., Faraday Trans. 1, 1977, 73, 871. 48 M. R. Powell and B. Rosenberg, Biopofymers, 1970, 9, 1403. 49 S. Bone, J. Eden and R. Pethig, Int. J . Quantum Chem., Quantum Biol. Symp., 1981, 8, 307. 50 M. Eigen and L. DeMaeyer, Proc. R . SOC. London, Ser. A, 1958, 247, 505. Paper 5/569; Receitled 3rd April, 1985

ISSN:0300-9599    

DOI:10.1039/F19868200143

出版商:RSC    

年代:1986

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1986, 82, 157-160 Aggregation in Aqueous Solution of the Dye Pyronine G John Gormally* and Susan Higson Department of Pure and Applied Chemistry, University of Salford, Salford A45 4 WT The aggregation in aqueous solution of the dye Pyronine G has been studied using spectrophotometry. A high-purity sample of Pyronine G was used, and it is shown that the dye spectra can be interpreted in terms of dimerization of the dye. Values of the dimerization constant at three temperatures are given. Pyronine G is a metachromatic xanthene dye with the structure shown in fig. 1. This dye has been used as a stain in microscopy,l as a fluorescent laser dye2 and in the spectrophotometric determination of platin~m.~ It has also been used in the study of dye-polyelectrolyte interactions4 and in a study of the kinetics of dye aggregation in s01ution.~ In spite of these activities, reliable data on the solution properties of this compound are rare6 owing to the low purity of most of the samples studied.For example, reported molar absorptivities for dilute aqueous solutions at 547 nm range from 1.2 x lo4 dm3 mol-1 cm-l [ref. (7)] to 10.7 x lo4 dm3 mol-1 cm-l [ref. (6)]. We present below some values of parameters which determine the spectrophotometric properties and association behaviour of Pyronine G obtained with a high purity sample of the dye. Experimental Commercially available samples of the dye (BDH Chemicals Ltd and Eastman Kodak Co.) were found to be very impure. A large proportion of the contaminating material was insoluble in methanol and could be removed easily.The dye extracted in this way had a molar absorptivity at 547 nm of 5.7 x lo4 dm3 mol-1 cm-l at a concentration of mol dm-3. Recrystallization from both methanol and ethanol was attempted, but was not successful owing to the high solubility of the dye in these solvents. The dye used in this study was synthesized using the method of Jacobsen et aL8 and had a molar absorptivity at 547 nm of 10.96 x lo4 dm3 mol-1 cm-l mol dm-3 dye in lov3 mol dmP3 HCl at 25 "C). This is very close to the value reported by Jacobsen et aL6 (10.7 x dm3 mol-1 cm-l) and suggests that our product is similar in purity to theirs. To our knowledge this is the highest value reported for Pyronine G and the purity of the substance prepared in this way is supported by microanalysis and n.m.r.investigations described in ref. (8). Absorbance measurements were made using a Pye Unicam SPS-300 spectrophoto- meter with thermostatting facilities. Sample cells had pathlengths of 40 mm, 10 mm and 2 mm and prior to use they were wetted with a 2% solution of dimethyldichlorosilane in trichloroethane (B.D.H. Chemicals Ltd) and then rinsed in distilled water. This treatment inhibits adsorption of dye onto the cell walls. Solutions were made up in mol dm-3 HCl to restrict the pH to ca. 3 as Pyronine G spectra are relatively insensitive to pH fluctuations in this r e g i ~ n . ~ ? ~ The presence of HCl also reduces the tendency of cationic dyes to adsorb onto the surfaces of glas~ware.~ Absorbances were measured at wavelengths from 500 to 570nm at 10nm intervals and at the a-band absorbance peak at 547 nm. The concentration range studied extended from (0.8 to 75) x mol dm-3.157158 d - I 5 52- "0 5 0 - 3 4 8 - 4 6 - 4 2 4 0 - Aggregation of Pyronine G I I I -6 - 5 - 4 4 4 + G * - $ - - - 100 " I P - 3 m E a I 0 - - 50 0 - 1 P 1 1 I I 450 500 A/ niii I 5 50 Fig. 1. The absorption spectrum of Pyronine G chloride (concentration 1.02 x lop5 mol dmp3 at 25 "C). The wavelengths at which absorbances were measured are indicated by vertical lines and examples of such data are given in fig. 2. - 6 4 Fig. 2. Plots of apparent absorptivity against dye concentration at ( a ) 550 and (b) 510 nm at a temperature of 30 "C. Data were obtained at nine different wavelengths and at three temperatures.Curvature of the plots indicates departure from the Beer-Lambert law and is attributed to dye aggregation.J . Gormally and S. Higson 159 Results Typical results showing the variation of apparent absorptivity with concentration are shown in fig. 2. Curvature in these graphs is characteristic of metachromatic dyes and is attributed to dye aggregation. At any wavelength, A, the apparent molar absorptivity can be written as \ n / where the E , are the absorptivities of species n (n = 1 for monomer, 2 for dimer, etc.) and A, are the corresponding concentrations. Atot is the total dye concentration and is n = A,+2K,A:+3K,K3AX+ . . . where K,, is the equilibrium constant for the association equilibrium (3) It is readily shown that if we consider aggregation only to the dimer stage that eqn (1) becomes Data such as are shown in fig.2 were fitted to this expression using the simplex methodlo and a minimum sum of squares criterion was used to find the values of E,, c2 and K , which gave the best fit. This was done for sets of data obtained at different wavelengths to check the wavelength independence of K,. The spread in values of K2 obtained is reflected in the error limits of the values given in table 1. Attempts to evaluate dimerization constants from data obtained at a single wavelength can lead to an optimistic assessment of the errors involved.ll The same procedure was adopted with data obtained at one wave- length, but at different temperatures. In this case consistency in the values of el and E , was sought.The values shown in table 1 indicate the consistency found. It was noticed that the value of the monomer molar absorptivity at 547 nm increased by ca. 0.3% per degree drop in temperature. This effect was observed in very dilute solutions in which the concentration of aggregates was immeasurably small, and it is too large to be accounted for by the increase in concentration which attends thermal contraction of the solution. A similar effect has been reported for solutions of xanthene dyes in ethano1.12 Aggregation was also considered to the trimer by retaining the third term in eqn (3) which then becomes a cubic in A,. This was solved for A, using the method in ref. (13) and the resulting expressions were used in conjunction with eqn (1) and (2) to determine &(A) in terms of E,,E,,E,,K, and K3.Unique values of K2 and K3 could not be found, indicating that whilst higher aggregates may exist in these solutions their concentration is too small to have a significant effect on our measurements. Conclusion Aggregation in aqueous solutions of Pyronine G within the concentration range studied can be adequately explained in terms of dimerization of the dye with the parameters given in table 1. In comparison with most other metachromatic dyes, Pyronine G is very highly absorbing in the visible region and has a relatively small dimerization ~0nstant.l~ From the values of Kz at different temperatures, estimates of various thermodynamic parameters can be obtained. Within the temperature range studied the values for AG and AH were found to be - 18.8 f 0.2 kJ mol-l and - 21 5 kJ rnol-l, respectively.The error in AH is large, but we consider this to be reasonable considering the very low concentration of dimer which exists in these solutions (for a dye concentration of mol dm-3 at 6 FAR 1160 Aggregation of Pyronine G Table 1. Examples of absorptivities and dimerization constants derived from an analysis of spectrophotometric dataa T "C E , (547) Cd547) Ed5 10) c2(5 10) ~ ~ / 1 0 3 25 1 1.4( 0.1) 0.7( 0.2) 4.3( * 0.1) 14.3 & 0.5) 2.0( * 0.1) _- ~ ~ ~ ~ _ _ _ _ 30 1 1.2( f 0.1) 0.75( & 0.2) 4.3( k 0.1) 14.5( k 0.5) 1.7( 40.1) 14.5 11.7( & 0.2) 0.95( 0.2) 4.4( k 0.1) 15.5( & 0.5) 2.7( 0.15) a Absorptivities (wavelength in parentheses) are measured in lo4 dm:3 mol-' cm-'.25 OC, the concentration of dimer is around 2 x lo-' mol dm-"). Further, we note that some of the reported values of thermodynamic parameters relating to dye aggregation imply a very optimistic assessment of accuracy as is evident from the wide range of values for the enthalpy and entropy of dimerization of acridine orange to be found in the literature [ref. (14)]. It seems possible that these quantities could be measured more accurately by use of the isoextraction method which Mukerjee and Ghosh applied to a study of aggregation in methylene blue.15 However, we feel that the values given above are realistic in that any more accurate determination, should this be required, would give results within the error limits specified. References 1 F. H. Kasten, Stain Technol., 1962, 37, 265. 2 B. I. Stepanov and A. N. Rubinov. Souiet Physics Uspekhi, 1968, 11, 304. 3 S. Jaya, T. P. Rao and T. V. Ramakrishna, Analyst (London), 1984, 109, 1405. 4 V. Vitagliano and L. Costantino, J . Phys. Chem., 1970, 74, 197. 5 W. Ohling, Ber. Bunsenges. Phys. Chem., 1984, 88, 109. 6 P. Jakobsen, H. Lyon and S. Treppendahl, Histochemistry, 1984, 81, 99. 7 K. Fujiki, C. Iwanaga and M. Koizumi, Bull. Chem. SOC. Jpn, 1962, 35, 185. 8 P. Jakobsen, A. P. Andersen, H. Lyon and S. Treppendahl, Microsr. Acta, 1983, 87, 41. 9 P. Mukerjee and A. K. Ghosh, J . Am. Chem. SOC., 1970, 92, 6403. 10 J. A. Nedler and R. Mead, Comput. J., 1965, 7, 308. 1 1 R. L. Reeves, M. S. Maggo and S. A. Harkaway, J . Phys. Chem., 1979, 83, 2359. 12 J. E. Selwyn and J. I. Steinfeld, J . Phys. Chem., 1972, 76, 762. 13 I. S. and E. S. Sokolnikoff, Higher Mathematics for Engineers und Physicists (McGraw-Hill, New York, 14 V. Vitagliano, in Aggregation Processes in Solution, ed. E. Wyn-Jones and J. Gormally (Elsevier, 15 P. Mukerjee and A. K. Ghosh, J . Am. Chem. SOC., 1970, 92, 6419. 1941), p. 86. Amsterdam, 1983), p. 276. Paper 51570; Received 3rd April, 1985

ISSN:0300-9599    

DOI:10.1039/F19868200157

出版商:RSC    

年代:1986

数据来源: RSC

摘要:

J. Chem. SOC., Furuduy Trans. 1, 1986,82, 161-165 Reactivity of the Hydrated Electron with (co-Carboxylatoalky1)ferricenium Zwitterions A Radiation Chemical Study G. Arthur Salmon Cookridge Radiation Research Centre, Cookridge Hospital, Leeds LSI 6 6Q B S. R. Logan* Department of Chemistry, University of Ulster, Coleraine, Northern Ireland BT52 1 SA Rate constants have been determined for the reaction of the hydrated electron with the substituted ferricenum zwitterions obtained by one-electron oxidation of 3-ferrocenylpropanoate, 4-ferrocenylbutanoate and 5-ferro- cenylpentanoate anions. These values are all close to 3 x 1Olo dm3 mol-l s-l and since, on the basis of the equation of von Smoluchowski, this implies that reaction occurs when the centres of the reactants are at a separation of at least 7 A, it would appear that these reactions are diffusion-controlled. Thus the sandwiching of the iron atom between two cyclopentadienyl groups does not appreciably alter the high reactivity towards the hydrated electron of the ferric part, whose complexes with inorganic Iigands are known to react at the diffusion-controlled limit.Previous work using ferrocenyl-substituted carboxylic acids explored the properties of the ferrocenyl group in aqueous solution with regard to its photochemical behaviourl and its reactivity2 with the hydroxyl radical. In both of these cases, oxidation of the ferrocenyl group takes place and a zwitterionic substituted ferricenium is obtained, with characteristic absorption peaks at ca. 255 and 625 nm, which is (in nearly all cases) quite stable and can readily be chemically reduced back to the corresponding ferrocenyl anion.It was considered of interest to determine how readily the hydrated electron reacted with these ferricenium species. One of the simplest ways of generating the ferricenium species is to y-irradiate an N,O-saturated solution of the ferrocenyl compound. However, since there is evidence2 that the reaction of the OH radical with a ferrocenylalkylcarboxylate is not exclusively an oxidation process, it would appear better to add bromide ion so that the oxidising species is converted into Br;: eLq + N,O -+ N, + o-("-+ OH) (1) (2) (3) OH + Br- --+ OH- + Br' Br' + Br- -+ Br; Br; + Fc(CH,),CO; -+ Fc+(CH,),CO; + 2 B r . (4) Thus the G-values of the ferricenium zwitterion would be equal to the sum of the G-values of the hydrated electrons and hydroxyl radicals scavenged at the prevailing concentrations of N,O and Br-.Thus it was arranged to carry out pulse-radiolysis studies on a range of samples of a ferrocenylalkylcarboxyl which had been oxidised to varying extents to the correspond- ing ferricenium zwitterion. The pseudo-first-order rate constant for decay of the hydrated electron should be a linear function of the concentration of the ferricenium species, and 161 6-2162 Reaction of the Hydrated Electron with Ferricenium Zwitterions allowing for the possibility of reaction with the ferrocenyl species, the gradient of this plot would be a measure of the difference of the second-order rate constants of the hydrated electron with the ferricenium and the ferrocenyl Epecies.The magnitude of the latter may be deduced from the intercept and thus the desired rate constant may be obtained. Experimental Materials The ferrocenyl-substituted carboxylic acids were prepared as previously describedl or were supplied by Prof. W. E. Watts. Other chemicals were of AnalaR grade and triply distilled water was used. Procedure All vessels used were first cleaned using permanganic acid followed by acidified hydrogen peroxide and then by numerous rinsings. Solutions of each of the ferrocenyl-substituted carboxylic acids were made up in 0.01 mol dm-" borax, which buffered them to pH 9. Sufficient KBr was also added to give a concentration of 0.02 mol drnp3. A sample of 30 cm3 of solution was taken and put into a vessel2 for y-irradiation which was fitted with a side-arm leading to a cylindrical quartz cell of 10 mm path length.The use of a septum cap and syringe needles permitted the solution to be deaerated and saturated with N,O. The spectrum of the solution over the range 200-800 nm was monitored periodically using a Pye Unicam SP8-100 spectrophotometer, in the course of y-irradiation with a 6oCo source of 10000 Ci nominal activity. It was found that the rate of growth of the characteristic ferricenium absorption peaks decreased slightly with time of y-irradiation. Samples were y-irradiated for different periods of time, to produce a suitable range of concentrations of the ferricenium species, After y-irradiation, each solution was thoroughly bubbled with argon over a glass frit in order to remove the dissolved N,O and it was then put into an all-glass syringe.The irradiation cell of the pulse-radiolysis system could be filled from these syringes through a flow system which allowed the sample to be changed without contamination with air. Immediately after the pulse-radiolysis measurements on each sample, a portion of the remaining solution was used to measure the absorbance of the solution at 628 nm in a 10 Inm cell, in order to estimate the concentration of the ferricenium zwitterion. Pulse Radiolysis The pulse-radiolysis system employed has been described elsewhere.4, Irradiation of the sample was by pulses of 3 MeV electrons generated by a Van de Graaff accelerator [High Voltage Engineering (Europa), KS-30001 and having a duration of either 25,50 or 100 ns.In experiments where the rate of reduction was being measured, the maximum initial concentration of the hydrated electron was 10 pmol dm-3. In order to minimise interference by the absorption of the Br; ion, the transient absorption due to the hydrated electron was monitored at 775 nm, using a silicon photodiode as detector. During the experiment a Chance OY6 filter was interposed between the xenon lamp and the irradiation cell to prevent photolysis of the sample prior to pulsing. Kinetic traces were recorded using a Tektronix 7912 transient digitizer and then fed to a computer which was used to evaluate the decay kinetics. At least six satisfactory traces were recorded for each solution used, with the irradiation cell being refilled with fresh solution after every second pulse.G.A . Salmon and S. R. Logan 163 3 - 2 - - I v1 r 0 F-, . 4 I , 0 0.1 0.2 0.3 A Fig. 1. Plot of k’, the pseudo-first-order rate constant for the decay of the hydrated electron, against A , the absorbance at 628 nm, from experiments with solutions of 5-ferrocenylpentanoate containing 0.02 mol dm-3 KBr which.had been y-irradiated to varying extents while saturated with N,O. [Fc(CH,),CO;] = 1.01 (0) and 3.07 (a) mmol dm-3. Determination of Ferricenium Concentration By means of the Fricke dosimeter,6 it was established that the mean dose rate in the 30 cm3 sample during y-irradiation was 59.8 Gy min-l. On this basis, from the initial growth of the absorbance at 628 nm when the solutions of the ferrocenylalkylcarboxylates were y-irradiated, values* of GE could be obtained.A solution containing 0.01 mol dmP3 borax and 0.02 mol dm-3 KBr and saturated with N,O by bubbling as already described, was irradiation with 100 ns pulses and the maximum absorbance at 360 nm due to the Br; ion was measured. The energy input to the irradiation cell under these conditions was calibrated using the oxygenated 0.01 mol dm-3 KI do~imeter,~ for which GE at 385 nm is known to be 3.18 x loP4 m2 J-l. Taking the extinction coefficient8 of Br; at 360 nm as 990 60 m2 mol-1 it was calculated that in the KBr solution under those conditions G(Br;) = (6.35+0.4) x lop7 mol J-l. Taking the G value of the ferricenium compound in the initial stages of y-irradiation as being equal to this figure, the extinction coefficient of the ferricenium species can then be obtained and thus the absorbances measured just after the pulse measurements can be converted into concentrations.Results From each kinetic trace recorded, the pseudo-first-order rate constant k’ for hydrated electron decay was calculated, typically over a period of between two and three half-lives. Taking the mean value of k’ for each solution, this rate constant was plotted against the absorbance measured at 628 nm just after the pulse-radiolysis measurements. The very small value of the intercept shows that the hydrated electron reacts only slowly with the ferrocenyl compound itself. A sample plot of k’ against absorbance is shown in figure 1 for experiments using 5-ferrocenylpentanoate.Two different concentrations of the latter were employed, but * Where G is the radiolytic yield of the light absorbing species in units of mol J-’ and E is its molar absorptivity in units m2 mol-’.164 Reaction of the Hydrated Electron with Ferricenium Zwitterions Table 1. Evaluation of the second-order rate constants gradient of k' cs. A GE/ lo-" k / 10"' dm" compound / 10' s-1 m2 J-' &/m2 mol-1 mol-1 s-l (2-carboxylatoethyl)ferricenium 9.66 k 0.35 2.1 1 k 0.07 33.2 & 3.2 3.2 k0.4 (3-carboxylatopropyl)ferricenium 9.60 k0.60 1.88 k0.07 29.6 k 3.0 2.8 & 0.5 (4-carboxylatobutyl)ferricenium 9.67 k0.20 2.11 k0.07 33.2 k 3.2 3.2 & 0.4 ~ ~~ ~ ~ ~ ~ ~ ~ _ _ _ _ _ _ _ _ _ _ _ ~ ~ _ _ _ - ~ ~ ~ ~ _ _ _ _ _ _ _ ~ - ~ ~ it may be seen that both sets of points fall on the same straight line.From each such plot, the least-squares line was obtained and the gradient measured. Table 1 lists for each compound the values obtained for the gradient of this plot and for GE. Since the irradiation cell as a path length of 10 mm, the former divided by E [and taking G as (6.35 & 0.4) x lop7 mol J-l] gives the second-order rate-constant for the reaction of the hydrated electron. No experiments were attempted with 2-ferrocenylethanoate, since it was known' that the corresponding dipolar ion was not stable over the necessary period of hours. In experiments with 3-ferrocenylacrylate it was found that the U.V. spectrum of the solutions changed appreciably between y-irradiation and the pulse experiments, so that it was not possible to assess what effect the presence of a double bond in the alkyl chain had on the rate of electron attack on the zwitterion.However, it was noted that the lifetime of the electron was unusually short in the unirradiated solution, indicating a much higher reactivity with this than with the other ferrocenyl compounds, and giving a rate constant of ca. 9 x lo9 dm3 mol-I s-l. Likewise, in experiments with 2-ferrocenylbenzoate it was noted that the absorbance of the ferricenium bands decreased by ca 20% between y-irradiation and the completion of the pulse-radiolysis measurements. If the data were to be treated in the same way as is described above, then the second-order rate constant for the reaction of the hydrated electron with the ferricenium zwitterion is found to be ca.2.5 x lolo dm3 mol-l s-l, on the basis of GE at 730 nm being 2.1 1 x lo+ m2 J-I. Also, the hydrated electron reacts rather faster with 2-ferrocenylbenzoate anion than it does with the co-carboxylalkyl- ferrocene anions, giving a rate constant of ca. 1.6 x lo9 dm" mol-1 s-l. In view of the ready autoxidation of the ferrocenylalkylcarboxylates and of the high reactivity of the oxidation product with the hydrated electron it is difficult to derive a reliable rate constant for the reaction of the latter with these anions, but it would appear that it does not exceed 0.5 x lo9 dm3 mo1-1 s-l. Discussion The relatively high rate constant for the reaction of the hydrated electron with 3-ferrocenylacrylate provides another illustration of the enhancing effect of a double bond in conjugation with an existing n-electron system.This mirrors the comparison of the rate constantss for the reaction of the electron with, on the one hand, benzene or toluene and on the other, styrene. With the (co-carboxylatoalky1)ferricenium zwitterions, the reactions of the hydrated electron are very fast. If we neglect the polar nature of these substrates and consider them as overall neutral species, then the rate constant for a diffusion-controlled reaction involving such species can be approximated by the equation'" of von Smoluchowski, k , = 4naD,,. Since the diffusion coefficient of the hydrated electron is known" to be 4.9 x lop5 cm2 s-l and that of the ferricenium species should be no more than 1 x cm2 s-l, this leads, on the basis of the rate constants listed in table 1, to a valueG.A . Salmon and S. R. Logan 165 of 0 of 6.9 A. This value is sufficiently large in relation to the dimensions of the ferrocene molecule that it would suggest that reaction does indeed take place on the first collision. This work represents the first measurement of the rate of reaction of the hydrated electron with a ferricenium species. In the context of the reactivity of eiQ with FeIII, the difficulty of preparing Fe3+ except in strongly acidic solutions necessarily restricts comparison to complexes of ferric ion. Studies with ions such as FeFi-, Fe(CN),3-, Fe(CN),N02- and Fe(EDTA)2- have shown1* that all of these reactions are extremely fast, in terms of the von Smoluchowski-Debye expression13 for reaction between ions. In the present work no ligand such as NO, which can itself react rapidly with eiq, is present, but clearly the cyclopentadienyl groups between which the iron is sandwiched do not provide a barrier to the transfer of the electron to the metal ion.Calculations14 of the charge distributions within the ferrocene and ferricenium species indicate that the effective charges on the iron atoms are + 1.39 and + 1.47, respectively. Hence, in the ferricenium complex the cyclopentadienyl rings each carry a nett charge of -0.24, as compared with their formal charges of - 1 .O. This effective electron deficiency of the C,H, rings suggest that they may act as effective bridges for the transfer of an electron to the metal centre. Thus the ferrocenyl/substituted ferricenium system is one in which both the oxidation of the former by the hydroxyl radical15 and the reduction of the latter by the hydrated electron take place at rates which are close to the diffusion-controlled limit.We thank C. Kilner for assistance with the experiments, Dr F. Wilkinson for developing the system for the evaluation of decay rates and Prof. W. E. Watts for providing samples of some of the compounds studied. We also acknowledge financial assistance from the S.E.R.C. for the provision of the computing facilities. References 1 E. K. Heaney and S. R. Logan, J . Cliem. Soc., Perkin Trans. 2, 1977, 1353; 1978. 590. 2 S. R. Logan and G. A. Salmon, J . Chem. Soc., Perkin Trans. 2, 1983, 1781. 3 E. K. Heaney, S. R. Logan and W. E. Watts, J . Organomet. Chem., 1978, 153, 229. 4 T. J. Kemp, J. R. Roberts, G. A. Salmon and G. F. Thompson. J. Phys. Chem., 1968, 72, 1464. 5 F. S. Dainton, E. A. Robinson and G. A. Salmon, J. Phqis. Chem., 1972, 76, 3897. 6 J. W. T. Spinks and R. J. Woods, An Introduction to Radiation Chemistry (Wiley. New York, 2nd edn, 7 G. V. Buxton, Proc. R. SOC. London, Ser. A, 1972, 9. 8 G, L. Hug, Optical Spectra of Nonmetallic Inorganic Transient Species in Aqueous Solution, NSRDS- 9 M. Anbar and P. Neta, Int. J. Appl. Radiat. Isot., 1967, 18, 493. 1976), p. 93. NBS 69 (U.S. Department of Commerce, Washington D.C. 1981), p. 52. 10 M. von Smoluchowski, Z. Phys. Chem., 1917,92, 129. 1 1 K. Schmidt and M. Anbar, J . Phys. Chem., 1969, 73, 2846. 12 M. Anbar, M. Bambenek and A. B. Ross, SelecteJSpecific Rates of Reactions oJ'Transients from Water in Aqueous Solution, I . Hydrated Electron, NSRDS-NBS 43 (U.S. Department of Commerce, Washington D.C. 1973). 13 P. Debye, Trans. Electrochem. Soc., 1942, 82, 265. 14 P. S. Bagus, U. I. Walgren and J. Almlof, J. Chem. Phys., 1976, 64, 2324. 15 G. A. Salmon and S. R. Logan, Radiat. Phys. Chem., 1984, 24, 593. Paper 51667 ; Receised 22nd April, 1985

ISSN:0300-9599    

DOI:10.1039/F19868200161

出版商:RSC    

年代:1986

数据来源: RSC

摘要:

J. Chem. Soc., Faraday Trans. 1, 1986, 82, 167-178 Radical Cations of Ethylene Oxide and other Oxiranes Prepared by Exposure to Ionizing Radiation An Electron Spin Resonance Study Jan Rideout, Martyn C. R. Symons* and Brendan W. Wren Department of Chemistry, The University, Leicester LEI 7RH Exposure of dilute solutions of tetramethyl oxirane in trichlorofluoro- methane to ,OC y-rays _-_ at 77 K gave a species identified as the ring-opened radical cation, (Me,~-O-~Me,)f. Its e.s.r. spectrum comprised 13 lines with A(lH) = 15.1 G. This species had an intense violet colour, and there was no evidence for the cyclic cation. Similarly, the cations of cyclopentene oxide and cyclohexene oxide gave red-violet colours and complex e.s.r. spectra which were interpreted in terms of the oxirane ring opening to give 6- and 7-membered cyclic cations only.Both the 2,3-cis-dimethyl and the monomethyl derivatives on exposure at 4 K gave poorly defined e.s.r. spectra and no coloration. On warming to ca. 30 K the spectra became better defined and at 77 K intense red colours appeared together with new e.s.r. features which could be analysed in terms of the open-ring structures. Oxirane itself gave a spectrum with splitting due to four equivalent protons, together with small extra splittings probably due to coupling to a chlorine nucleus. The proton coupling of ca. 16 G is close to that for the two a-protons of the open-chain radical (cRMzOCHMe)+, which strongly supports an open structure. However, there was no visible absorption band under all conditions in CFCl,.In marked contrast, a very similar e.s.r. spectrum of five lines was obtained from solutions in SF, at 77 K, but these solutions had an intense orange-red coloration. These results are tentatively interpreted in terms of a cyclic cation for oxirane in CFCl,, but an ' open ' cation in SF,, both having very similar proton hyperfine coupling constants. Similarly, the acquisition of colour and change in the e.s.r. spectrum of the monomethyl and dimethyl derivatives is explained in terms of ring-opening in the 4&70 K region for solutions in CFCl,. ~~ It is now well established that exposure of dilute solutions of many neutral compounds in solvents such as the freons, especially CFCl,, to ionizing radiation results in the formation of the corresponding radical cati0ns.l Sometimes, these undergo various types of rearrangement prior to detection, and sometimes they form weak complexes with solvent molecules, but, generally, good spectra of the primary cations are obtained.The first report of the e.s.r. spectrum of the oxirane cation showed that this comprised a well defined quintet due to hyperfine coupling to four equivalent protons.2 The small magnitude of this coupling (16 G) clearly established that the structure differed from the normal n structure found for ether cation^^?^ and, in accord with a recent theoretical treatment,5 it was suggested that the SOMO (semi-occupied molecular orbital) was the alternative a-orbital (a,). This cation was studied independently by Snow et aL6 who suggested -__. that the C-C bond had, in fact, completely broken, to give the open structure [H,~-O-CH,]+, the unpaired electron being largely confined to the two methylene carbon atoms.That the open structure is more stable than the cyclic structures is supported by recent calc~lations,~-~ and we stress that there need not be a large barrier to ring-opening.* Strong support for the open structure came from optical studies of 167168 E.S.R. of Oxirane Radical Cations oxirane derivative cations;1° a good case being made that a band in the 560 nm region was due to the ‘open’ cations, the closed species having no visible absorption. One of the species studied was the tetramethyl derivative (2max = 560 nm)Io and, in collaboration with Williams and coworkers, it has been clearly shown by e.s.r.spec- troscopy that this species has indeed undergone ring-opening.ll Thus the case for ring-opening is strong. The aim of the present work was to examine this case in depth and, in particular, to see if there was any chance of detecting the closed-ring cations by suitable choice of solvent and temperature. Experiment a1 Dilute solutions (ca. 0.1 % mole fraction) of various oxiranes in CFCI, or CCl, were cooled in liquid nitrogen to form polycrystalline masses. The samples were exposed to WO y-rays up to a dose of 0.5 Mrad, under liquid nitrogen. For work at 4 K, X-rays were generated using an Andrex NDT Products (UK) Gemini 160 kV constant potential (X-ray generating) unit fitted with a beryllium window and water cooler. Ethylene oxide in SF, was prepared by mixing known pressures of the two gases giving ca.1 : 1000 v/v and fast freezing the mixture into a tube which was irradiated under the same conditions as the CFCI, and CC1, samples. The oxiranes were prepared, where necessary, by oxidation of the corresponding olefin with m-chloroperbenzoic acid. The resultant oxiranes were purified by preparative g.1.c. using a 20% P.E.G. column. The proton n.m.r. spectra of the final oxirane products were recorded to confirm their identification and the absence of impurities. E.s.r. spectra were recorded on a Varian E-109 X-band spectrometer, using an Oxford Instruments 4 K insert, where necessary. The samples were pulse annealed by decanting the coolant and recooling to 77 K whenever significant changes in the e.s.r.features were detected. A Varian variable temperature assessory was also used to obtain e.s.r. spectra at specific temperatures above 77 K. Results As stressed above, the e.s.r. and optical results for the tetramethyl derivative cation are clear and so we start with a description thereof. Similarly, the two cyclic derivatives give ‘open’ cations (I and 11) unambiguously. We then discuss the 2,3-dimethyl derivative, for which two different species were detected. Finally, we turn to the parent oxirane cation and its monomethyl derivative and consider their structures in the light of the foregoing results, eb 0 11 Tetramethyl Oxirane Radical Cations The e.s.r. spectrum for this cation obtained at 77 K comprises a set of 13 narrow lines with A(lH) = 15.2 G, which is almost certainly due to the open ring cation (fig.l), (Me2CoCMe2)+, having equal spin-density on the two carbon atoms. It is noteworthy that the isotropic splitting is greater than that for the corresponding allylic cations (ca. 14.7 G),12 as expected because of the positive-charge effect,13 but this increase is relatively small. The irradiated samples had an intense pink colour. The eye is a good spectrometer for absorption bands in the centre of the visible region and this colour proves the presence of a band in the 550 nm region. This accords well with the optical study of Bally et al.loJ . Rideout, M . C. R. Symons and B. W. Wrcn 169 + P25G I * H 10G , Fig. 1. First-derivative X-band e.s.r. spectrum for tetramethyl oxirane in CFCl,.after exposure to 6 ” C ~ y-rays at 77 K, showing features assigned to the ‘open’ cation (Me,COCMe,)+.[The + 6 features were detectable at high gain.] We conclude that the tetramethyl derivative forms an open cation directly at 77 K and that this is characterised by a methyl proton coupling of ca. 15 G and by an intense pink colour. This result does not give us a value for the a-proton coupling of a conclusively open oxirane cation so we turned our attention to the two cyclic oxirane cations, (I and 11). Cyclopentene Oxide and Cyclohexene Oxide Cations The e.s.r. spectrum of the cyclopentene derivative (I) at ca. 130 K was well defined (fig. 2). The analysis indicated in the figure gave A(2H) = 9 G, A(2H) = 15.3 G and A(2H) = 31 G. Clearly, the two strongly coupled protons must be in the axial sites for an open cation, but it is not directly obvious which of the two smaller coupling constants belongs to the two a-protons. Both of the alternative assignments are possible.That the species is indeed the ‘open’ cation (I) is clearly established by the appearance of a bright pink colour. Results for the cyclohexene derivative (11) were more difficult to interpret, as can be judged by the spectrum shown in fig. 2(b). In fact, the spectra were remarkably temperature sensitive and it is clear that a range of conformational changes must be taking place. Despite these complexities, however, we can conclude from the similar colour and overall width of the e.s.r. spectrum that the ring-opened cation was formed. In view of the ambiguity in assignment for the a-proton coupling mentioned above, we turned to the monomethyl derivative.The results for this species, described below, were still ambiguous, hence we synthesised the 2,3-dimethyl derivative. cis-2,3-Dimethyl Oxirane Cations The initial spectrum obtained at 77 K was complex and clearly due to two species [fig. 3(a)]. The narrow set of lines (a) were lost irreversibly on annealing, leaving, at cu. 90 K a set of nine lines, with a splitting of ca. 17 G [fig. 3(b)]. This changed reversibly to a set of eight lines on cooling to 77 K [fig. 3 ( c ) ] . The sample had a strong pink colour throughout.170 E.S.R. of Oxirane Radical Cations 1 3240 G Fig. 2. First-derivative X-band e.s.r. spectra for ( a ) cyclopentene oxide and (h) cyclohexene oxide in CFCl, after exposure to V o y-rays at 77 K and annealing to a.150 K. showing features assigned to the ‘open’ cations. [(b) shows the selective line-broadening discussed in the text.] Weconclude that the 8-9 line speciesis the ‘open’cation and that A(aH) = AWH) = 16- 17 G. This result is definitive in establishing that the a-proton coupling is in the 16 G range and hence the 9 G splitting obtained from the cyclopentene derivative is due to the two equatorial hydrogen atoms. Given an average P-proton coupling of 15-16 G, values of 31 and 9 G for axial and equatorial protons are well accommodated by the cos20 law for the axial and equatorial protons. These results appear to confirm the ‘open’ assignment for the parent oxirane cation. However, they leave open the assignment of the extra features in fig.3(u). Exposure of the dimethyl derivative at 4 K gave a colourless solution with very poorly defined e.s.r. features which clearly differed from those of the ‘open’ cation. On warming to 77 K, the pink colour grew in intensity together with the features shown in fig. 3(a).J . Rideout, M. C. R. Symons and B. W. Wren 1 32356 171 + 3 32355 10G - H - \ Fig. 3. (a) and (b). For legend see next page. The Oxirane Cation Further studies of this cation confirm previous results using CFCI, as solvent. The spectrum comprises a set of five broad lines with weak extra splitting assigned to coupling to a single chlorine nucleus.6 This was lost reversibly on annealing. The most striking aspect of these results is the complete absence of any red colour.More concentrated solutions acquired a pale yellow colour whilst still exhibiting the normal five-line spectrum. This leaves us with two clear possibilities: (i) the ‘open’ cation is formed, as suggested by the splitting of 16 G and the colour is only found for alkyl-substituted cations; or172 E.S.R. of Oxirane Radical Cations I W H 1QC , 1 32356 Fig. 3. First-derivative X-band e.s.r. spectra for dimethyl oxirane in CFC1, after exposure to 6oCo ?-rays at 77 K, ( a ) showing features assigned to the open cation (p) [see (c)] together with features (a) indicated in the stick diagram, which were lost irreversibly on annealing to ca. 120 K; ( 6 ) showing the nine-line spectrum assigned to the open cation at ca. 150 K, with freely rotating methyl groups and ( c ) showing how this changes reversibly to an eight-line spectrum with the same total width on cooling to 77 K.(ii) this colourless cation has the closed ring 2 A , structure originally proposed.2 We stress that, for equal concentrations, all the previous ' open ' cations have exhibited intense pink colorations, in accord with the results of Bally et a1.l0 In contrast, the pale yellow colour found for the parent cation accords well with the optical spectrum given for the cation derived from (111), which only had a band in the U.V. region.'" In the hope of resolving this problem, we decided to study the oxirane cation in SF,, which is a far less constricting solvent than CFCl,, in the expectation that this might encourage ring-opening. CFCl, molecule is much smaller than an SF, molecule.Thus the increase in size on ring opening is sterically opposed in CFCl,, effectively increasing the barrier to ring opening in this solvent). The e.s.r. results were almost identical with those for solutions in CFCl,, apart from the absence of extra splitting from chlorine and a slightly smaller hyperfine coupling (fig. 4). However, the samples had an intense orange-red colour after irradiation, suggesting an absorption band at ca. 500 nm. Finally, we turn to our results for the monomethyl derivative. 2-Methyl Oxirane Cations The e.s.r. spectra obtained for solutions of this cation in CFCI, proved to be difficult to interpret (see below). However, there can be little doubt that the spectrum obtained on annealing to ca. 130 K [fig.5(a)] is due to the 'open' cation, having the expected red colour. Exposure at 4 K gave rise to very broad e.s.r. features even at minimum microwave powers. However, the samples were colourless. On annealing to 25-30 K a better defined spectrum was obtained [fig. 5(b)] which was quite different from that assigned to the 'open' cation [fig. 5(a)]. After annealing to 77 K features assigned to theJ . Rideout, M . C. R. Symons and B. W. Wren I + H 1 10G , 3205G I73 Fig. 4. First-derivative X-band e.s.r. spectrum for oxirane in SF, after exposure to T o prays at 77 K, showing features assigned to the open (orange-red) cation. 'open' cation grew in and the sample acquired an intense red colour. This colour change was irreversible. The spectrum shown in fig.5(h) can be analysed in several ways since it is so poorly defined. However, we feel confident that it cannot be due to the 'open' cation. Discussion There has been a spate of theoretical studies related to the structure of both closed and open forms of the oxirane cation.5+ It seems to be generally agreed that the first-formed cation should have the normal n-structure, with a 2Bl ground state (IV). However, the first excited state, (V) can stabilise by stretching the C-C bond and this is the structure favoured by Mollere and Houk5 and suggested by ourselves2 for the radical cation in CFCl,. If, as found by Clark,7 the ,A, species in its most favourable relaxed form has a very long carbon-carbon bond (1.78 A) and effectively planar carbon centres, an isotropic hyperfine coupling of ca.16 G is reasonable. *A 1 2B 1 However, there is now general agreement that the most stable form is the open structure, though there has been considerable disagreement as to the degree of allylic stabilisation experienced by this structure. Originally, it was thought that an asymmetric structure with one long and one short C-0 bond (VI) was most stable but, most recently, a near-planar, C,, structure seems to have been accepted (VII>.7-9 The former is, of course, ruled out by the e.s.r. results, although they do not prove that the favoured structure is necessarily the limiting planar species (VII) and there may well be some degree of twist, especially for the substituted cations. H174 E.S.R. of Oxirane Radical Cations i 3240G 10 G -H 1 326% Fig.5. First-derivative X-band e.s.r. spectra for 2-methyl oxirane in CFC1, ( a ) after exposure to 6oCo prays at 77 K and annealing to ca. 140 K, showing features assigned to the open cation and (b) after irradiating at ca. 4 K and warming to ca. 40 K, showing the sextet tentatively assigned to the colourless parent cation. Again, a hyperfine coupling of 16 G is very reasonable for such a structure and it is generally assumed that the e.s.r. results establish that ring-opening does, indeed, occur at 77 K for the oxirane cation. In our view, the oxirane cation formed in CFCl, is not the ‘open’ cation, despite the fact that the lH hyperfine coupling is close to that found for the ‘open’ cation in SF, and for the a-proteins of the ‘open’ dimethyl derivative.Our arguments are the following: (i) there is no visible absorption band for the species in CFCl,, whereas that in SF, has an intense orange-red colour (ymax in the 500 nm range). Furthermore, careful annealing of the CFC1, solutions failed to yield a coloured cation prior to an irreversible e.s.r. spectral change discussed below. (ii) The mono- and di-methyl derivatives gave colourless cations on irradiation at 4 K and intense red colours grew in irreversibly onJ. Rideout, M . C. R. Symons and B. W. Wren 175 annealing, together with marked e.s.r. spectral changes. The initial e.s.r. spectra cannot be assigned to the primary 2B, state (IV), but could well be due to the ‘ A , state (V). Thus, in these cases, the alternative structures do not give rise to identical e.s.r.spectra. (iii) Two species were formed from the 2,3-dimethyl derivative at 77 K, one of which was lost irreversibly at ca. 100 K. It seems probable that the unstable species was the 2Al-state cyclic cation. (iv) Weak hyperfine splitting assigned to a chlorine nucleus appears in the spectrum for the colourless oxirane cation in CFCl,. This is a surprising result for either of the proposed structures since common experience is that such an interaction is only observed when a well defined a-bond can be formed between a highly localised radical cation centre and chlorine. However, from the reported correlation between hyperfine coupling and ionization potential,14. l5 we predict a maximum splitting of ca. 50 G which is far greater than that observed (ca.4 G). For the cyclic cation, this reduction suggests that the interaction is between oxygen and chlorine as in (VTII), but that the SOMO remains largely confined to the stretched C-C a-bond. However, for the ‘open’ cation, we suggest that any tendency to interact at one side would distort the radical, making the SOMO favour the remote side, the limiting structure being that shown in (IX). This would, of course, be formed directly from the ‘orthogonal’ structure (V). Just such a structure has been invoked16 to explain why the SOMO is effectively confined to one -CH2 unit in the ‘open’ cation derived from the cyclopropane cation,17 rather than being equally distributed between both outer methylene groups, as in the present case. I I CL-CFCI, v I11 IX It is noteworthy that a dependence of ring-opening on solvent has recently been detected for the cyclopropane cation, which remains cyclic in CFCl,, but which undergoes ring-opening in the less constricting solvent, CFC1,-CFCI.l i In the light of these considerations, we conclude that, whilst ring-opening is clearly thermodynamically favoured, there is a significant barrier for ring-opening from the ,A, state and that this is inhibited by a constricting solvent. Spectral Interpretation Several aspects of the detailed interpretation of the e.s.r. spectra have been ignored in the above discussion in the interests of simplicity. In some cases, we are unable to offer more than these general statements, but, in others, some further details have been forthcoming. 2-Methyl Oxirane Cations We had expected that the ‘open’ cation would have proton hyperfine coupling constants similar to those for the parent cation, the cis-2,3-dimethyl derivative and the tetramethyl derivative.However, we cannot fit the spectra on this basis and have been forced to consider an asymmetric structure in which the electron is biased in favour of the -CH, unit and the positive charge is biased in favour of the -CHMe unit. A reasonable fit for the spectrum at ca. 130 K is obtained if the a-protons of the CH, unit have a = 12 G, whilst the a-proton coupling for the CHMe group is 21 G and the methyl proton splitting is 21 G. [These values were obtained from CFC1, solutions: the features were better resolved for solutions in CCl,, but the hyperfine coupling constants were slightly smaller (table l).] We suggest that this asymmetry is induced by the stabilising effect of176 E.S.R.of Oxirane Radical Cations Table 1. E.s.r. data for some oxirane radical cations structure lH hyperfine coupling/Ga ~ _ _ ~~~ oxirane of cation solvent T/K 0 vMe 0 \/ 0 Me-Me Me, - Me, 0 \/ 8 0 c) 0 closed CFCl, open SFti closed CFCI, open CFC1, CCl, closed CFC1, open CFCl, open CFCl, open CFC1, open CFCI, 150 77 I 77 4 - - 4 130 77 77 77 16.3 15.7 b 4H ca. 12, 2H ca. 21 4H ca. 12, 2H CLI. 19 b 8H CLI. 16.5 12H. 15.2 2H, 9, 2H, 15.3, 2H, 31 b colour colourless orange-red colourless red colourless violet violet violet violet ~ ~~ ~~~~ a G = lo-, T; No unambiguous interpretation available. hyperconjugation from the methyl group, which is expected to stabilise a carbocation centre more than a radical centre.In contrast, the SOMO’s,for corresponding unsym- metrical neutral ally1 radicals remain almost evenly distributed between the terminal atoms. Unfortunately, the spectra at very low temperatures for the colourless samples are too broad to warrant any attempts at interpretation. The six-line spectrum obtained in the 3G50 K region may be due to the primary cation, but unfortunately, it is difficult to analyse in terms of reasonable expectation for the 2A, state cation. The results (six lines with splitting ca. 10.5 G) appear to require three protons with A z 10.5 G and one with A = 21 G. It is difficult to reconcile these results with that for the cyclic oxirane cation unless the spin distribution is again highly asymmetric. 2,3-Dimethyl Oxirane Cations Features tentatively assigned to the cyclic cation were never sufficiently defined for proper analysis.This problem was exacerbated by their marked temperature dependence. However, as indicated in fig. 3(a) the extra features obtained at 77 K can be interpreted in terms of two protons with A = 16.5 G and two others with A = 21.6 G. There are, however, several extra features of lesser intensity which may belong to this species. If the two protons with A = 16.5 G are identified as the two ‘a’ protons, then the sum of the methyl proton coupling constants equals 43.2 G which gives 7.2 G per proton. Hindered rotation of both groups could well lead to a set of septets, with the M , = 3 and 0 lines narrow and the remainder broad.We tentatively assign these results to the cyclic 2 A , state cations. The reduction in methyl C-H hyperconjugation may be a resultJ . Rideout, M . C. R. Symons and B. W. Wren 1 32306 177 I - 1 Fig. 6. First-derivative X-band e.s.r. spectrum for oxirane in CFC1, after exposure to ‘j0Co prays at 77 K and annealing to ca. 150 K, showing the triplet species discussed in the text. of the weak a-bonding or of a certain degree of residual pyramidality at each carbon atom. Changes in the spectra of the cyclic derivatives were highly complex on temperature variation. Since the overall spectral widths remained constant and the changes were reversible, it is clear that changes in conformation or equilibria between different conformations are responsible rather than chemical changes.Since these are not our present concern, we have abandoned attempts at detailed interpretation. Other Irreversible Spectral Changes Only one well defined change occurred on annealing towards the melting point of the medium (ca. 160 K). That was for the parent oxirane cation, shown in fig. 6. This new spectrum, obtained irreversibly in all experiments, comprises a 1 : 2: 1 triplet with A(lH) z 20 G. We offer two explanations. One is that a bimolecular reaction between the cation and another oxirane molecule occurs to give (X), followed by rapid ring opening of the intermediate radical to give (XI). [The parameters shown in this insert are taken from ref. (1 8).] Radical (XI) could certainly be responsible for the triplet, but there was no sign of the cyclic radical (X) and we would be surprised if this were to undergo rapid ring-opening at ca.150 K, judging from the liquid-phase stability of this species.’* The alternative is that the 2A, state cyclic cation undergoes ring-opening, but that this species reacts efficiently with a solvent molecule to give the adduct (IX). This would also be expected to exhibit a triplet spectrum of the type observed. (6C) H , 0 H2C-C 4 ‘ti X XI We thank Professor F. Williams for sending us his results for tetramethyl oxirane cations, Professor Haselbach for interesting discussions and the S.E.R.C. for a grant.E.S.R. of Oxirane Radical Cations References 1 M. C. R. Symons, Chem. SOC. Rev., 1984, 13, 393. 2 M. C. R. Symons and B. W. Wren, Tetrahedron Lett., 1983, 24, 2315. 3 H. Kubodera, T. Shida and K. Shimokoshi, J . Phys. Chem., 1981, 85, 2583. 4 M. C. R. Symons and B. W. Wren, J . Chem. SOC., Perkin Trans. -3, 1984, 51 1. 5 P. D. Mollere and K . N. Houk, J . Am. Chem. SOC., 1977, 99, 3226. 6 L. D. Snow, J. T. Wang and F. Williams, Chem. Phys. Lett., 1983, 100, 193. 7 W. J. Bouma, D. Poppinger, S. Saebs, J. K. MacLeod and L. Radom, Chem. Phys. Lett., 1984, 104, 8 T. Clark, J. Chem. Soc., Chem. Commun., 1984, 666. 9 D. Feller, E. R. Davidson and W. J. Borden, J . Am. Chem. SOC., 1984. 106, 2513. 198. 10 T. Bally, S. Nitsche and E. Haselbach, Heh. Chim. Acta, 1984, 67, 86. 11 F. Williams, Faraday Discuss. Chem. Soc., 1984, 78, 57. 12 J. K. Kochi and P. J. Krusic, J . Am. Chem. SOC., 1968,90.7157; H. Elson. S. W. Mao and J. K. Kochl, J . Am. Chem. SOC., 1975,97, 335. 13 J. Brivati, R. Hulme and M. C. R. Symons, Proc. Chem. Soc., 1961, 384; R. Hulme and M. C. R. Symons, J . Chem. Soc., 1965, 1120. 14 A. Hasegawa, J. Rideout, G. W. Eastland and M. C. R. Symons, J . Chem. Res. (S), 1983, 258. 15 T. Clark, A. Hasegawa and M. C. R. Symons, Chem. Phys. Lett., 1985, 116, 79. 16 M. C. R. Symons, Chem. Phys. Lett., 1985, 117, 381. 17 X-Z. Qin and F. Williams, Chem. Phys. Lett., 1984, 112, 79. 18 A. J. Dobbs, B. C. Gilbert and R. 0. C. Norman, J . Chem. SOC. ( A ) , 1971, 124; G. Behrens and D. Schulte-Frohlinde, Angew. Chem., 1973, 85, 993; H. Itzel and H. Fischer, Helr. Chim. Acta, 1976, 59, 880. Paper 51680; Receiced 24th April, 1985

ISSN:0300-9599    

DOI:10.1039/F19868200167

出版商:RSC    

年代:1986

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1986,82, 179-188 Application of Microelectrodes to the Study of the Li1Li-i Couple in Ether Solvents Part 2.-Temperature Dependence William M. Hedges and Derek Pletcher Department of Chemistry, The University, Southampton SO9 5NH Cyclic voltammetry and potential step experiments at microelectrodes have been used to probe the influence of temperature (over the range 253-333 K) on the deposition and dissolution of lithium metal in tetrahydrofuran containing lithium hexafluoroarsonate. The temperature strongly influences both the kinetics of nucleation of the lithium phase on other conducting materials (Ni, Cu, C) and of the Li I Li+ couple after the new phase is formed. The exchange current density for the LilLi+ couple at 293 K is found to be 4 mA cm-2 for a solution containing 0.84 mol dmP3 Li+, a much higher value than previously reported in the literature, suggesting that these earlier data were adversely affected by experimental problems, surface films and/or IR drop, The energy of activation for the couple was determined to be 56 kJ mol-'. The temperature dependence of the LiAll Li+ couple is also discussed.The past twenty years have seen many groups active in the development of lithium batteries and, indeed, several types have now reached the stage where they are commercially available. One group of rechargeable lithium batteries employs an electrolyte consisting of a lithium salt dissolved in an aprotic organic medium, preferably an ether or a mixture of ethers. Our understanding of the physical chemistry of these systems, however, lags considerably behind their exploitation. Experimental problems, e.g.uncompensated IR drops, water in the electrolyte and films on the lithium surface, make it difficult to obtain reliable and reproducible data. As a result, much of our knowledge of the deposition and dissolution of lithium metal and of the kinetics of the Li I Li+ couple in ether solvents is inferred from galvanostatic charge cycling experiments. Such experiments have been reported in several solvent compositions and electrolytes and the influence of additives has been disc~ssed.~-~~ Another approach has been to use propylene carbonate as a model solvent and studies in this solvent include charge cycling,l2> l3 a.c. impedance,14 cyclic voltammetryl53 l6 and e1lips0metry.l~ Even with this solvent, with a high dielectric constant, the experimental data are clearly not free from distortion due to IR drop at moderate current densities, but when such studies have been extended to ethers, gross distortion is evident.Takei18 reported that lithium deposition could not be observed by cyclic voltammetry in tetrahydrofuran while Matsuda et al. found very high Tafel slopes, (500 mV)-l, even at low current densities.lg Recently, however, Genders et a1.20 have demonstrated that the use of microelectrodes allows potential step and sweep techniques to be used with ether solvents without the data being distorted by IR drop. The LiAl I Li+ couple is also of interest for application in battery technology.21 The formation of the LiAl alloy and its ability to cycle charge has been widely studied, propylene carbonate again being the most widely used ~ o l v e n t .~ ~ - ~ ~ While charge cycling efficiency has been the most common method, a.c. 25 potential step 179180 Li I Lif Couple in Ether Solvents techniques 25 and cyclic voltammetry1j. l6 have been used in the study of this couple. Once more, IR drop problems will distort the data at higher current densities in ether solvents. In this paper we discuss the temperature dependence of the kinetics of both the Li I Li+ and the LiAl I Li+ couples in tetrahydrofuran containing lithium hexafluoroarsonate ; microelectrodes are again used to obtain reliable and undistorted experimental data. 28-30 Experimental Tetrahydrofuran (THF) (East Anglia Chemicals) was refluxed under an N, atmosphere and above equimolar amounts of freshly cut sodium and benzophenone and then distilled immediately prior to experiments. The blue benzophenone anion radical serves to scavenge oxygen, water and peroxides from the ether.The electrolyte was always electrochemical grade lithium hexafluoroarsonate (USS Agri-Chemicals) and it was used without purification. Solutions were prepared inside an Ar-circulating atmosphere dry box, but most experiments were carried out in a closed cell outside the box. Solutions were deoxygenated with oxygen-free Ar (B.O.C. Ltd). All experiments were carried out in two electrode cells. The working electrodes were prepared from wires, copper (radius 40 pm), nickel (25 pm) and aluminium (25 pm) all supplied by Goodfellow Metals or carbon fibre (radius 8 pm) from Harwell.They were sealed into soda glass using a vacuum pump to collapse the glass onto the wire and the glass was then cut so that only a disc at the end of the wire was exposed to solution. Electrical contact was either via a thicker nichrome wire joined to the far end of the microwire to the disc or via a lead plug (formed by melting lead pellets) to a thicker nichrome wire. The micro-disc electrodes and the glass surrounds were initially ground on a wheel and then polished with various grades of alumina powders on a moist polishing cloth. Between experiments the electrode surface was polished with the finest grade (0.05 pm) alumina powder. The other electrode was prepared by pressing lithium foil (Foote Mineral Co.) on to a Pt wire contact and combined the functions of counter and reference electrode. The cells were undivided with a volume of 10 cm3 and hence the potentials of the working electrode were measured versus the equilibrium Li I Li+ potential in the cell electrolyte, i.e.as an overpotential directly. The resistances of the cells were of the order of 5-10 kQ. Hence with most experiments the maximum IR drop was ca. 10 mV despite the high current densities 0.1-200 mA cm-,. Water in the electrolyte solution led to a cyclic voltammogram with a reduction wave 400 mV positive to lithium deposition and this wave serves as a convenient monitor for contamination of the solutions by water. The potential of the working electrode was programmed using a Hi-Tek function generator, model PPR1.The currents passing through the cell were amplified using a home-built amplifier whose output was displayed either directly on a Farnell xy recorder or stored in a Gould digital oscilloscope type OS4100 and then displayed on a recorder. When electrical noise was a problem, the experiments were carried out in a Faraday cage. Results The Li I Li+ Couple Fig. 1 shows cyclic voltammograms for a nickel electrode (radius 25 pm) run in THF/LiAsF, (0.83 mol dm-3) at 255 and 295 K. At 295 K, fig. 1 (a), the I-Ecurve is very similar to that reported earlier for a copper electrode.20 On the forward sweep, the current density is very low until a potential is reached where nucleation of lithium metal can occur on the nickel surface.Then the current density increases rapidly (and can reach a value above 0.2 A cm-2) and continues to increase after the direction of scan is reversed as the reduction of Li+ is now occurring at lithium centres whose surface area is increasing.W. M . Hedges and D. Pletcher 181 /I 10.2 PA or 10 mA cm-2 V Fig. 1. Cyclic voltammograms for an Ni electrode ( r = 25 pm) in THF containing LiAsF, (0.83 mol dm-3). v = 100 mV s-l. (a) 295 K; (b) 255 K. Deposition continues until 0.0 V us. the Li I Li+ reference electrode, i.e. the equilibrium potential for the Li I Li+ couple in the electrolyte, but on going to positive potentials, the Li metal is oxidised and a stripping peak is observed. The stripping efficiency is close to 100%. The behaviour at Ni differs from Cu only in that the overpotential for nucleation of the lithium is larger; at 100 mV s-l nucleation occurs at - 140 mV on Ni and -90 mV on Cu.The cyclic voltammogram on a carbon fibre is also similar, here nucleation commencing at -270 mV at 100 mV s-l. The I-E curve at 255 K, fig. 1 (b), immediately demonstrates that both the deposition and corrosion of Li require higher overpotentials at this lower temperature and certainly the current densities are lower at all potentials. In principle, such changes could be due to changes in diffusion coefficient, the kinetics of nucleation and/or the kinetics of the Li I Li+ couple. but the distinction is difficult to make from cyclic voltammograms. The deposition and corrosion of lithium were both explored further using the single and double potential step experiments described previously.2o Fig.2 ( a ) and (b) show deposition and stripping transients recorded at 253 K and using a Ni electrode. As previously found for a Cu electrode at room temperature, the transients at negative potentials have the classical shape for a nucleation and growth process. Initially the current is almost zero, but this increases with time until a steady state value is reached; the timescale of the transient decreases and the plateau current increases as the overpotential is made more negative. The stripping transients also have the expected shape. The lithium corrodes at almost constant rate until the layer on the Ni surface becomes depleted and then the current decreases to zero. The initial steady state rate of corrosion is also a strong function of the overpotential. Similar experiments were carried out at 273,293 and 313 K and in all cases the qualitative shapes of the transients were the same although at any overpotential, the magnitude of the current densities increases and the timescales of the transients decrease as the temperature is increased.At 253 and 293 K, and at overpotentials of - 175 and -200 mV, the shapes of the early rising parts of the transients were analysed in more detail. The best fit between the182 Li I Li+ Couple in Ether Solcents ( A ) 70 1s 5 0 10 t / s -220 mV 0 2.5 t l s 5 .O Fig. 2. I-t transients for an Ni electrode in THF/LiAsF, (0.84 mol dm-3) at 253 K. Potential sequences: ( A ) 0.0 V to (a) - 165, (b) - 175, (c) -200 and (d) -210 mV; (B) 0.0 V to -220 mV for 12 s then to (a) 300, (b) 350, (c) 400 and (d) 450 mV.experimental data and various I us. tn functions was for I us. t2 and the corresponding I4 us. t plots are shown in fig. 3. The analysis of the slopes of these plots, given in table 2, is discussed later. The plateau currents from the transients for the deposition and corrosion of lithium metal should be capable of analysis in terms of the kinetics of the Li I Li+ couple. It was found in our earlier work20 that the value of the exchange current density, I,, for the Li I Li+ couple at room temperature is sufficiently high that when the steady state I us. ul data are plotted as log I vs. q, it is not possible to identify convincing linear Tafel regions. On the other hand, a plot of I us.7 clearly shows an inflection close to q = 0. HenceW. M . Hedges and D. Pletcher 183 0.3 0.6 0 4 8 tl s t/ s Fig. 3. I:-t plots for data taken early in the I-t response following a step from 0 V to (u) - 175 and (6) -200 mV. ( A ) 293 K; (B) 253 K. Ni electrode in THF/LiAsF, (0.67 mol dm-"). it was considered worthwhile to analyse the steady state data using the Hickling equation31 which is a rearrangement of the full Butler-Volmer equation and does not only apply when the back reaction may be neglected. Hence it allows data closer to 7 = 0 to be included and also allows the information from the oxidation and reduction to be included on the same line. Steady state data were measured at four temperatures. At each temperature data for both oxidation and reduction were obtained from a series of experiments designed to ensure that the data correspond to a clean Li surface with the minimum influence from surface films.The polished Ni microelectrode was placed in the cell at 0.0 V and its potential was first stepped to a value where a thin layer of Li (typically 1 C cmP2) was deposited (the potential employed varied with T ) . Then the potential was stepped to a value in the range 250 to -250 mV and the current for either oxidation or reduction was recorded as soon as it reached a steady state value (usually within a few s). Such data plotted according to eqn (1) are shown in fig. 4. It can be seen that acceptable straight lines with the expected slopes are obtained, although inevitably over a restricted range of overpotentials because of the limitations from mass transport control at higher overpotentials.The values of I, obtained from the intercepts and the values of the transfer coefficient (or) obtained from the slope are reported in table 1 and the former were used to construct the Arrhenius plot in fig. 5. The value of the activation energy obtained is 56 & 5 kJ mol-l. A similar procedure to that described above was used to investigate the variation of I , with the concentration of Li+ in solution. Over the range 0.02-2.0 mol dm-3 (but without added base electrolyte) a plot of I, us. [Li+]i is linear, confirming that the reduction is first order in lithium ion.184 LiILi+ Couple in Ether Solzwits -200 -100 . q/mV us. Li I Li+ Fig. 4.Plot of steady-state I-7 data according to eqn (1). The electrode is electroplated Li on Ni in THF/LiAsF, (0.84 mol drnp3). (a) 254, (b) 273, (c) 293 and (d) 313 K. Table 1. Kinetic data for the Li I Li+ couple taken from plots in fig. 4 T/K IJmA cmP2 a 254 0.16 0.50 273 0.94 0.57 293 4.0 0.55 313 26 0.49W. M . Hedges and D. Pletcher 185 7 L N 1 - c 3 0 d E P . M - 0 - - 1 - 1 - 7 --I- - --- 3 0 3 . 2 3.4 3.6 3 8 4 0 1 0 3 KIT Fig. 5. Variation of I, with Tpresented as an Arrhenius plot. The LiAl I Li+ Couple The electrochemistry of lithium at an aluminium electrode will differ from that at Ni, Cu or C because lithium and aluminium form well defined alloy phases; a-LiAl contains 0-7% Li while P-LiAl contains 47-56% Li. Hence the electrochemistry will involve diffusion of lithium in the aluminium lattice and the equilibrium potential for the electrode process will be shifted positive by the free energy of dissolution of lithium in the aluminium. It has been reported25 that the equilibrium potential for a P-LiAlILi+ couple in propylene carbonate is 38 1 mV us.Li I Li+. Fig. 6 shows cyclic voltammograms for an A1 electrode in THF/LiAsF, (0.56 mol drnp3) at 294 and 333 K. There are several significant differences from the results at Ni or Cu. At the lower temperature, the forward sweep shows no current until a potential of - 60 mV when the current density rises steeply. On reversal of the potential scan, there is a large nucleation loop and reduction continues until 330 mV when the current becomes anodic and a rather broad oxidation peak is observed.Several features of the curve should be noted: (a) the current density reaches almost 0.1 A cmP2 at the negative limit showing the high rate of diffusion of Li in the A1 lattice (if the current density is allowed to go significantly higher, Li metal forms on the surface and a second anodic peak is observed at about 0.3 V); (b) the stripping efficiency for Li is only ca. 70% as not all of the lithium diffuses back to the surface of the Al; ( c ) the shape of the reverse scan around I = 0 clearly indicates that the kinetics of the LiAlILi+ couple are significantly slower than those for Li 1 Li+; ( d ) the broad shape of the anodic stripping peak probably results from the need for diffusion of Li through the LiAl phase to the surface of the electrode; (e) the potential for zero current on the reverse scan may not correspond to the equilibrium potential because the solid LiAl phases may not reach an equilibrium composition or lattice structure on the short timescale of these experiments (the changes in potential would be expected as the system tends to the steady state would lead to a more positive value); and (f) the nucleation potential is at least 390 mV, much higher than observed on Ni or Cu.186 Li I Lis Couple in Ether Soluents 25 mA cm-2 /I] I l i 1.4 0 0 -0.15 + b ' V us Ll1 LI+ 1 1.6 0.0 -0.075 +E/V us Li I Li+ Fig.6. Cyclic voltammograms for an A1 electrode (v = 25 pm) in THF/LiAsF, (0.56 mol dmP3). v = 100 mV s-l. (a) 294 K ; (h) 333 K. On increasing the temperature to 333 K, the cyclic voltammogram shows higher current densities, a lower nucleation overpotential (nucleation now occurs positive to 0.0 V us. Li I Li+) and the slope of the I-Ecurve around zero current is steeper, confirming an enhancement of the kinetics of the LiAl I Li+ couple.Potential step experiments were also carried out at room temperature. Steps to negative potentials again led to well defined rising I-t transients although their reproducibility was not as good as at Ni or Cu. Such transients would confirm that, unlike amalgam formation, the first stage in the production of LiAl is a nucleation process and then growth of the small centres of the new phase. Our data are entirely consistent with the model proposed by Owen and Ma~ke11,~~ where formation of LiAl occurs as three dimensional centres on the surface of the aluminium, although we cannot rule out the possibility that the nucleation phenomena arise because of the need to remove an Alto, film before the LiAl can form.Discussion The experiments reported in this paper illustrate well the advantages of microelectrodes for studies in poorly conducting media. Reliable and readily interpretable data can be obtained from simple experiments and even in a solvent such as THF the data are free of experimental artefacts due to IR drop or charging currents. Certainly, we believe that this approach can give a detailed understanding of the electrochemistry of reactive metals and a good insight into battery systems based upon them.W. M . Hedges and D. Pletcher 187 Table 2.Analysis of the potential dependence of data taken from I-t transientsa - 175 mV -200 mV 253 K 293 K 253 K 293 K -~ ~ __ ~ ~~ slopes of 18 us. t 0.68 21 1.56 83 plots/ 1 0-4 mAh cm-l s-l.. . ratio of rate constants The exchange current densities reported in table 1 are higher than those reported in the literature1> l4 for ether solvents, but that at 293 K is a similar value to those reported for Li I Li+ in propylene ~arb0nate.l~ In addition, the transfer coefficient was always close to the expected value of 0.5 and we believe that much lower values reported in the literature are due to the influence of uncorrected I R drops. The energy of activation determined for the Li 1 Li+ couple, 56 5 kJ mo1-1 compares with values33 of 28 kJ molt1 for AglAg+ in non-complexing media (e.g.NO;) and 50-80 kJ mol-1 in complexing media (e.g. CN-), in line with the expectation that Li+ is strongly solvated in ether solvents. The linear 14-t plots obtained from the analysis of the lower rising portion of the I-t transients at - 175 and -200 mV at 253 and 293 K would be compatible with a deposition mechanism whereby instantaneous nucleation is followed by three-dimensional growth of the Li phase under kinetic control. Moreover the total shape of the transient is consistent with this interpretation if overlap of the centres is followed by linear expansion of the Li layer still under kinetic control. The shape of the initial part of the transient is then given by Is = ( p2 ) kgt 2F7cN,M2 4 where M is the atomic weight of Li, p its density, N , the total number of nucleation sites at the applied potential and k is the growth rate constant.If the latter is considered to be the rate constant for the electron transfer to Li+ at the applied potential and No is initially considered to be independent of potential and temperature, it is possible to test the total shape of the transients. The ratio of rate constants for the reduction of Li+ at - 175 and - 200 mV may be calculated in three ways: (i) from the slopes of the I+-t plots, (ii) from the ratio of the steady state plateau currents and (iii) from the Tafel equation; such a comparison is shown in table 2. It is found that there is reasonable agreement between the values from the three ways of estimating the ratio. In fact, the ratio of rate constants estimated from the limiting currents is slightly low because the reaction at - 200 mV is probably showing a small measure of diffusion control.The ratio determined from the I$-t plots is also close although the value of No at the more negative potential might be expected to be higher. The experiments with an aluminium electrode show clearly that LiAl is formed by a188 LilLi+ Couple in Ether Solvents phase formation process and that the lithium both goes into and leaves the A1 lattice at a surprisingly high rate if a sufficient overpotential is used; the kinetics of the LiAl I Lit couple are slow compared with the Li I Li+ couple, a factor which will influence the high rate capability of any battery with an LiAl electrode. References 1 Lithium Batteries, ed.J-P. Gabano (Academic Press, London, 1983). 2 A. N. Dey, Thin Solid Films, 1977, 43, 131. 3 V. R. Koch, J. Power Sources, 1981, 6, 357. 4 N. A. Hampson, M. Hughes and S. A. G. R. Karunathilaka, J . Power SourcrJ, 1984, 12, 83. 5 V. R. Koch and J. H. Young, J . Electrochem. Soc., 1978, 125, 1371. 6 J. L. Goldman, R. M. Mank, J. H. Young and V. R. Koch, J . Electrochem. Sue., 1980, 127, 1461. 7 K. M. Abraham, J. L. Goldman and D. L. Natwig, J . Electrochem. Soc., 1982, 129, 2404. 8 V. R. Koch, J. L. Goldman, C. J. Mattos and M. Mulvaney, J . Eleclrochem. Soc., 1982, 129, 1. 9 K. M. Abraham, J. S. Foos and J. L. Goldman, J . Electrochem. Soc., 1984, 131, 2197. 10 P. G. Glugla, J. Electrochem. Soc., 1983, 130, 113. I 1 J. S. Foos and J. McVeigh, J . Electrochem.SOC., 1983, 130, 628. 12 R. D. Rauh and S. B. Brummer, Elecfrochim. Ac~u, 1977, 22, 75. 13 R. Selim and P. Bro, J . Electrochem. Soc., 1974, 121, 1457. 14 R. V. Moshtev, J. Power Sources, 1984, 11, 93. 15 J. 0. Bessenhard, J . Electrounal. Chem., 1978, 94, 77. 16 E. J. Frazer, J . Electroanal. Chem., 1981, 121, 329. 17 F. Schwager, Y. Geronov and R. H. Muller, J . Electrochem. Soc., 1985, 132, 285. 18 T. Takei, J. Appl. Electrochem., 1979, 9, 587. 19 Y. Matsuda, M. Morita and K. Takata, J . Electrochem. Soc., 1984. 131, 1991. 20 J. D. Genders, W. M. Hedges and D. Pletcher, J . Chem. Soc., Furadaj. Trans. I , 1984, 80, 3399. 21 W. R. Fawcett, Soaiet Electrochem., 1983, 19, 1044. 22 A. N. Dey, J . Electrochem. SOC., 1971, 118, 1547. 23 B. M. Rao, W. R. Francis and H. A. Christopher, J. Electrochem. Soc., 1977, 124, 1490. 24 I. Epelboin, M. Froment, M. Garreau, J. Thevenin andD. Warin. J . Electrochem. Soc., 1980,127,2100. 25 A. S. Baranski and W. R. Fawcett, J . Electrochem. Sue., 1982, 129, 901. 26 T. R. Jow and C. C. Liang, J . Electrochem. Soc., 1982, 129, 1429. 27 A. S. Baranski, W. R. Fawcett, T. Krogulec and M. Drogowska, J . Electrochem. Soc., 1984,131, 1750. 28 R. M. Wightman, Anal. Chem., 1981, 53, 1125A. 29 J. 0. Howell and R. M. Wightman, Anal. Chem., 1984, 56. 524. 30 A. M. Bond, M. Fleischmann and J. Robinson, J . Electrounal. Chem., 1984, 168, 299; 1984, 172, 11. 31 P. A. Allen and A. Hickling, Trans. Faradaji Soc., 1957, 53, 1626. 32 J. R. Owen and W. C. Maskell, J. Electrochem. Soc., 1985, 132, 1602. 33 R. Tamamushi, Kinetic Parameters of Electrode Reaction& of Metallic Compounds (IUPAC - Butter- worth Press. Tonbridge, 1975). Puper 5/797; Receiced 13th Ma?,, 1985

ISSN:0300-9599    

DOI:10.1039/F19868200179

出版商:RSC    

年代:1986

数据来源: RSC