摘要:

J . Chem. Soc., Faraday Trans. I, 1985, 81, 207-213 Enthalpies of Mixed-micelle Formation BY MICHAEL J. HEY* AND JOHN W. MACTAGCART? Department of Chemistry, The University, Nottingham NG7 2RD AND COLIN H. ROCHESTER Department of Chemistry, The University, Dundee DDl 4HN Received 4th May, 1984 Critical micelle concentrations have been measured for binary mixtures of an anionic and a polyoxyethylenated non-ionic surfactant in aqueous solution. The data were used to calculate the molecular interaction parameter p in the Rubingh theory of mixed-micelle formation. Enthalpies of mixed-micelle formation determined calorimetrically were then compared with theoretical values based on 8. Differences found between expected and observed enthalpies are attributed to deficiencies in the theory with respect to the temperature dependence of p and the variation of counterion binding as the micelle composition changes.Mixtures of polyoxyethylene-type non-ionic and anionic surfactants in aqueous solution are known to form mixed micelles.' A synergistic effect is observed in that the critical micelle concentrations (c.m.c.) in the mixtures are lower than those of the pure components when the latter are approximately the same.2* The stabilization of the mixed micelles relative to pure micelles has been interpreted in terms of a thermodynamic model by Rubingh4 in which non-ideal mixing in the micelle takes place with a negative enthalpy change. Starting from the pseudo-phase-separation description of micellization the Rubingh theory allows predictions of the c.m.c.to be made as a function of the overall composition of the mixture. These are in good agreement with experiment if a molecular interaction constant (p) is treated as an adjustable parameter. The justification for the theory is based on the observation that a single p value serves to represent the data remarkably well over the complete range of mixing for a particular system. Recently the model has been generalized for multicomponent systems and comparisons of c.m.c. measured in ternary systems with predicted values show good agreement whenPvalues are used which were independently determined from binary systems.j Since the excess enthalpy of mixing in the micelle is related to p in the Rubingh theory it is of interest to determine this experimentally for a system of known p so that the validity of the theory can be established.To do this we have taken sodium dodecyl sulphate (SDS) and 0-n-octyltetraethylene glycol (C,E,), which have similar c.m.c., and from mixed c.m.c. values have used the theory to calculate p. Calorimetric measurements were then made to obtain the enthalpy change for formation of mixed micelles and the results compared with the values from the theory. EXPERIMENTAL MATE RIALS All surfactant solutions were made up immediately before use in ion-exchanged water that had been triply distilled, the first distillation being from alkaline potassium permanganate f Present address : Unilever Research, Port Sunlight Laboratory. Quarry Road East, Bebington, Wirral, Merseyside L63 3JW. 207208 HEATS OF MICELLIZATION solution.Distillations were carried out under an atmosphere of nitrogen, and the distillate, which was collected in a leached Pyrex vessel, was used within 12 h. SDS (B.D.H., specially purified for biochemical work) was used as supplied and stored in a vacuum desiccator. A sample of C,E, with a specified purity of 98.6% and a cloud point of 43.2 "C (1 % solution) was supplied by the Port Sunlight Laboratory of Unilever Research. Before use it was stored frozen at - 18 "C under nitrogen. SURFACE TENSIONS The surface-tension measurements were made by the Wilhelmy-plate technique, using glass microscope cover slips which had been roughened by grinding with silicon carbide + aluminium oxide paste. The plate was suspended from the arm of an electronic microbalance which could be raised and lowered using a vernier screw gauge.Solutions were contained by a double-walled glass vessel thermostatted at 25 "C. The complete apparatus was enclosed in a wooden case with plastic doors and set up in a constant-temperature room with a dust-controlled atmosphere. Prior to use all the glassware was cleaned with a mixture of nitric and hydrofluoric acids (5% ) then thoroughly rinsed with distilled water. CALORIMETRY Heats of dilution of the surfactant solutions were measured using an LKB 8700-2 reaction calorimeter in which the reaction vessel was immersed in a thermostatted bath maintained at 25 & 0.05 "C. Data were analysed according to the Regnault-Pfaundler method using the computer curve-fitting technique of BorrelL6 A test of the accuracy of the calorimeter was made by measuring the enthalpy of solution of tris(hydroxymethy1)aminomethane in 0.1 mol dmP3 aqueous hydrochloric acid.Two determinations gave AH = -29.78 and -29.72 kJ mol-l, which agree well with the literature' value (- 29.74 kJ mol-l). In each calorimetric experiment enthalpies of dilution and demicellization were deter- mined by breaking an ampoule containing the appropriate volume of 0.5 mol dmP3 total surfactant solution into 25 em3 of distilled water to give final concentrations in the range (1-20) x loP3 mol dm-3. A correction was made for the small heat change (0.10t0.02 J) associated with the breaking of the ampoules. RESULTS AND DISCUSSION The c.m.c. of the individual surfactants and their binary mixtures were determined by finding the points of discontinuity in plots of surface tension against the logarithm of the concentration.Fig. 1 shows the results as a function of a, the mole fraction of SDS. A 'vat-shaped' dependence was obtained in agreement with published data for the same system.2 To fit the Rubingh theory to the data the c.m.c. values were substituted into the equation x 2 In (c, a/c, x) = (1 - x ) ~ In [em( 1 - a)/c2( 1 - x)] where x is the mole fraction of SDS in the micelle at the c.m.c., c, is the c.m.c. of the mixed system and c, and c2 are the c.m.c. of pure SDS and C,E,, respectively. Eqn (1) was solved iteratively for x and the result substituted into eqn (2) to give a (1) An average value of -3.0 was obtained in this way over the composition range of the mixtures.This was then used to calculate theoretical c.m.c. values from the equations (3) (4)M. J. HEY, J. W. MACTAGGART AND C. H. ROCHESTER 209 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 Fig. 1. Critical micelle concentration plotted against overall mole fraction of SDS. The continuous line shows the fit of Rubingh theory with p = -3.0. cy -20 Fig. 2. Heat of dilution plotted against the reciprocal of the final concentration for different values of a : A, 0; 0, 0.1 ; 0, 0.4; a, 0.6; V, 0.9 and V, 1 (arrows indicate critical micelle concentrations). The line in fig. 1 shows that the experimental data are reasonably well fitted by the theory. Fig. 2 shows the results of the calorimetric experiments. AHdi1 is the heat change per mole of total surfactant produced by diluting solutions from 0.5 mol dm-3 to final concentrations in the region of the c.m.c.Differences in AHdil over the range of final concentrations will include contributions from the dilution of both monomers and micelles together with demicellization heats. However, at the low concentrations obtaining in the neighbourhood of the c.m.c. the dilution heats will be very small and may safely be neglected in comparison with the demicellization effect. The constancy O f AHdil at concentrations below the c.m.c. for pure SDS and pure C,E, confirms that this approximation is valid at least for the monomers. Micellization enthalpies are related to dilution enthalpies by wherefis the fraction of surfactant which is present in micellar form.Sincefis related to the total concentration of surfactant c by 8 FAR 1210 HEATS OF MICELLIZATION 1 .o 0.8 0.6 X 0.4 0.2 0 / 0.2 0.4 0.6 0.8 1.0 (Y Fig. 3. Micellar composition at the c.m.c. plotted against overall composition : broken line, p = 0; continuous line, p = -3.0. the micellization enthalpy can be obtained from It can be seen from the slopes of the lines in fig. 2 that surfactant mixtures with a 2 0.6 have very small micellization heats, like SDS alone.8 For lower SDS contents micellization heats become endothermic with AH,,, = 13.5 If: 1.5 kJ mo1-1 for pure C,E4 micelles. This result is of comparable magnitude to micellization heats quoted in the literature for the similar compounds C,E, (15.5f4 kJ m01-l)~ and C,E, (20.1 f 0.8 kJ mol-l).lo Before interpreting heats of mixed-micelle formation the micellar composition variable x has to be related to a by using the following equation from the pseudo-phase-separation model : where f l and f i are the activity coefficients for the two components in the mixed micelle.Eqn (4), (5) and (8) were used to calculate x as a function of a for p = 0 (ideal mixing) and p = - 3. The curves are shown in fig. 3. According to the Rubingh theory the excess molar enthalpy of micellization is given by HE = x( 1 -X)/?RT (9) and since we can write for the theoretical enthalpies of micellization in the SDS/C,E4 system AHmic = xAH,i,(SDS) + (1 - X) AHmic(C,E,)+ HE AHmic/kJ mol-* = 13.5(1 -x)-33RTx(l -x). (10) (1 1) Finally, fig. 3 allows the enthalpies to be expressed in terms of a.M.J. HEY, J. W. MACTAGGART AND C. H. ROCHESTER 21 1 14 1 - -2 L Fig. 4. Heat of micellization plotted against overall mole fraction of SDS. Points show experimental results, broken line predicted values for /3 = 0, continuous line predicted values for /3 = - 3.0. A comparison of the theoretical and measured heats is given in fig. 4. It is clear that a discrepancy exists, with the measured heats being more exothermic than eqn (1 1) predicts. In order to explain the difference we suggest that the regular solution assumption made in the Rubingh theory should be modified as described below. The simplest class of non-ideal binary mixtures is that for which the composition dependence of the excess chemical potentials of the components is given by p p = w( 1 - X ) 2 (12) and p? = wx2 (13) where w is taken to be independent of temperature.17 Since pF = RT lnfi the p parameter of Rubingh is related to w by p = w/RT and we obtain G" = HE = WX(~ - x ) = ~ ( l -x)PRT (16) A more general theory allows w to be temperature-dependent, in which case dP = -x(l -x)RP- d T and SE is no longer zero but is given by dw d T SE = -x(I - x ) - = - ~ ( 1 -x) (19) 8-2212 HEATS OF MICELLIZATION No data have been reported for /3 values in mixed micelles at different temperatures, but Rosen and Zhaoll have applied Rubingh’s approach to mixed-monolayer formation at the aqueous-solution/air interface and have obtained surface /3 para- meters for C,,E, and sodium dodecyl sulphonate at 298 and 3 13 K.The temperature coefficient for their monolayer p value over this range was 0.027 I C l , indicating a decrease in the non-ideality of mixing at higher temperatures.If dp/dT is put equal to this for the SDS/C,E, micellar system in eqn (18) the micellization enthalpy (20) becomes AHmi,/kJ mol-l = 13.5(1 -x)-20x(1 -x) in place of eqn (1 1) and the predicted heats now become closer to the measured heats. However, lack of knowledge of the proper value for da/dT renders quantitative comparisons unprofitable. It is sufficient to conclude that a positive temperature coefficient for a, of the order of magnitude found for the C,,E,/sodium dodecyl sulphonate monolayer, could bring theory and experiment into agreement. A second effect missing from the original theory arises from the variation in the degree of ionic dissociation as the composition of the micelle alters.It has been established that for sodium dodecyl sulphonate/C,,E, and SDS/C,,E, l2 micelles there is a decrease in the dissociation of the ionic component as the ionic/non-ionic ratio increases. Meguro et aZ.13 have suggested that binding of counterions is prevented by steric hindrance from the polyoxyethylene chains. Osmotic and electrical measure- ments show that the degree of dissociation falls fairly rapidly as the ionic mole fraction increases to ca. 0.4 and then more slowly as the pure ionic micelle is approached. In the case of SDS-containing micelles the final degree of dissociation is close to 0.2.149 l5 The enthalpy change caused by the binding of Na+ ions to dodecyl sulphate micelles has been measured calorimetrically16 as - 4.6 kJ mol-l.In order to predict the enthalpy change for mixed-micelle formation therefore we need to include an endothermic term AHdis(a) which is a function of a. The ideal enthalpy of mixing can then be written as with an upper limit for AHdis(a) as 01 approaches zero of 3.7 kJ mol-l. Thus the effect of counterion binding is in the opposite direction to the error caused by assuming w to be independent of temperature. The latter, however, is likely to be greater in magnitude, with the net result that measured heats are more exothermic than predicted by the Rubingh theory. Further tests and refinements of the model can only be made when more data are available for both the temperature dependence of p and the extent of counterion binding.Nevertheless, the present work clearly demonstrates that a regular-solution approximation which is more general than that adopted by Rubingh is needed to explain enthalpies of mixed micellization. We wish to thank the S.E.R.C. for the award of a CASE studentship to J. W.M. and Unilever Research for making available the facilities of the Port Sunlight Laboratory where the surface-tension measurements were made. J. M. Corkill, J. F. Goodman and J. R. Tate, Trans. Furuduy Soc., 1964, 60, 986. N. Nishikido, J . Colloid Interface Sci., 1977, 60, 242. D. N. Rubingh, in Solution Chemistry of Surfactants, ed. K. L. Mittal (Plenum Press, New York, 1979), vol. 1, p. 337. P. M. Holland and D. N. Rubingh, J . Phys. Chem., 1983, 87, 1984. P. Borrell, Thermochim. Acta, 1974, 9, 89. * H. Lange and K-H. Beck, Kolloid Z . Z. Polym., 1973, 251, 424.M. J. HEY, J. W. MACTAGGART AND C. H. ROCHESTER 213 ’ J. 0. Hill, E. Ojelund and I. Wadso, J. Chem. Thermodyn., 1969, 1, 1 1 1. a K. S. Birdi, Colloid Polym. Sci., 1983, 261, 45. lo J. M. Corkill, J. F. Goodman and J. R. Tate, Trans. Faraday Soc., 1964, 60, 996. l 1 M. J. Rosen and F. Zhao, J. Colloid Interface Sci., 1983, 95, 443. l 2 F. Tokiwa and N. Moriyama, J. Colloid Interface Sci., 1969, 30, 338. l 3 K. Meguro, H. Akasu and M. Ueno, J. Am. Oil. Chem. Soc., 1976,53, 145. C. Botre, V. L. Crescenzi and A. Mele, J. Phys. Chem., 1959, 63, 650. L. Shedlovsky, C. W. Jakob and M. B. Epstein, J. Phys. Chem., 1963, 67, 2075. J. S. Rowlinson and F. L. Swinton, Liquids and Liquid Mixtures (Butterworths, London, 3rd edn, 1982), chap. 5. J. M. Corkill, J. F. Goodman and S. P. Harrold, Trans. Faraday Soc., 1964, 60, 202. l 8 D. J. Eatough and S. J. Rehfeld, Thermochim. Acta, 1971, 2, 443. (PAPER 4/7 16)

ISSN:0300-9599    

DOI:10.1039/F19858100207

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J . Chem. SOC., Furaduy Trans. 1, 1985, 81, 215-221 Infrared Spectroscopic Study of the Adsorption of Hydrogen and Carbon Monoxide on Highly Dehydroxylated Thoria BY JEAN LAMOTTE AND JEAN-CLAUDE LAVALLEY* Laboratoire de Spectrochimie, ERA 824, Institut des Sciences de la Matiere et du Rayonnement, Universite de Caen, 14032 Caen Cedex, France AND VINCENZO LORENZELLI Laboratorio di Chimica, Facolta di Ingegneria, Fiera del Mare Pad. D, Universita di Genova, 16129 Genova, Italy AND EDOUARD FREUND Institut Frangais du Petrole, B.P. 31 1, 92506 Rueil Malmaison Cedex, France Received 9th May, 1984 Hydrogen adsorption on thoria activated at 973 K (Tho2-973) leads to weak bands at 11 15 and 860 cm-l and a shoulder at 910 cm-' due to partially reversibly adsorbed bridged species Th-H-Th.The uptake is very small. The adsorption of CO on Tho2-973 turns the disc orange-yellow and gives rise to a large number of species, some of which are reversibly adsorbed. These species are characterized by a band at 2171 cm-' and are coordinated to Lewis-acid sites. In addition to the carbonate and formate species, at least four types of irreversibly adsorbed species are formed. Three of them are characterized by bands in the 1200-1000 cm-l frequency range and can be assigned to v(C0) vibrations of -0-C-C-O- species; another possibility is the formation of Cog- species because of the strong basicity of thoria. The irreversibly adsorbed species are very sensitive to 0, and H,O but are not reduced by H,. Irreversibly adsorbed hydrogen species react with reversibly adsorbed CO species or CO gas, leading to species similar to those formed on adsorption of HCHO.Copper on thorium oxide (thoria) has been found to be a very active catalyst in the synthesis of methanol from CO+H,.l Moreover, Tho, itself is active in this reaction,'* the main product being (CH,),O. Therefore it will be interesting to study, using infrared spectroscopy, the mechanism of the CO + H, interaction and the nature of the active sites of Tho, in a similar way as for ZnO., No studies have been devoted to H, adsorption and the CO + H, interaction on Tho,. It has been reported that CO adsorption4 leads to formate species and to another species characterized by an i.r. band at 1290 cm-l. Moreover a band at 2143 cm-l corresponds to weakly adsorbed CO.* Using e.p.r. spectroscopy it has been found that CO adsorption forms radicals assigned to a negatively charged CO ad~orbate.~ we report here our results for the adsorption of H, and CO on Tho,.We concentrate our attention on highly dehydroxylated samples in order to compare these results with those recently obtained for the adsorption of H,7 and COs on alkaline-earth-metal oxides. After studying the acid-base properties of 215216 100, H, AND CO ADSORPTION ON Tho, A - - ? I 3700 2700 1500 1100 700 wavenum ber/cm-’ Fig. l.I.r. spectra of hydrogen (-) and deuterium ( . . . .) adsorbed on Tho,; P, = 66 kN m-, (the background has been subtracted). EXPERIMENTAL The Tho, (Rhbne Poulenc) used had a specific area of 120 m2 g-l. Further details are given in ref.(6). The 30 mg discs were activated at 973 K under vacuum (denoted Th0,-973). Spectra were recorded at room temperature using a FTIR-MX-1 Nicolet spectrometer. In the present paper we report the spectra obtained after subtraction of background absorbance. The gravimetric measurements were carried out in a conventional McBain thermobalance. High-purity hydrogen, deuterium and carbon monoxide (Air Liquide) gases were passed through a liquid-nitrogen trap before introduction into the sample cell. RESULTS The i.r. spectrum of Tho,-973 did not contain any bands due to carbonate species6 Because of the high degree of light scattering it was very difficult to study the spectra at wavenumbers > 3300 cm-l. H, ADSORPTION As noted for Mg0,7 hydrogen uptake on Th0,-973 is very small, almost at the limit of sensitivity of the volumetric apparatus.When the equilibrium pressure, P,, is lo2 N m-2, the amount of hydrogen adsorbed at 298 K was ca. 2.3 pmol g-l, e.g. No infrared bands due to H, adsorption appeared between 1200 and 800 cm-I on Th0,-770. However, the introduction of a large amount of H, (P, = 66 kN m-,) on Th0,-973 created two broad bands at 1 1 15 and 860 cm-l and a shoulder at 910 cm-l (fig. 1). Outgassing decreased their intensity, especially that of the I 1 15 cm-l band. In the 3800-3000 cm-l range we noted the appearance of a Th-OH band due to a free hydroxyl group at 3660 cm-l (sharp) with a shoulder at 3640 cm-l. These are less sensitive to outgassing and therefore could not be related to the broad bands between 1300-700 cm-l.These sharp bands are weak and could arise from homolytic dissociation of H, in very small quantities. When D, was used instead of H,, corresponding v(0D) bands appeared at 2700 (sharp) and 2685 cm-’ (shoulder), while in the 1200-700 cm-l range one broad band appeared at 800 cm-l and its intensity decreased on outgassing. molecule nm-2, which is in agreement with the value for Mg0.7 co ADSORPTION Several species were formed on the adsorption of CO on Tho,-973, while the white disc turned orange-yellow. The reversible species are characterized by a band atJ. LAMOTTE, J-C. LAVALLEY, V. LORENZELLI AND E. FREUND 21 7 . - - - - _ _ . . . . :.-. . . . . . . . . , , 2T h v 90 I: C .I m .I 5 a * I 781 L ...... 2200 2000 1600 14'00 1200 1600 wavenum ber/cm - 1 Fig.2. 1.r. spectra of irreversibly adsorbed CO species on Tho,: (-) room temperature; ( - . * a ) 403 K and (---) 523 K (the background has been subtracted). 21 71 cm-l with a shoulder at 2160 cm-l. Quantitative measurements showed that ca. 70 pmol g-l (P, = 26 kN m-,) are reversibly adsorbed. The irreversible species lead to many bands showing their multiplicity (fig. 2). We distinguish three main types. ( a ) A band at 2100cm-l. (b) Two complex and weak absorptions in the ranges 1600-1500 and 1400-1300 cm-l. The two broad bands (centred near 1520 and 1320 cm-l) are probably due to small amounts of surface carbonates and the sharp bands to formate species [v,,(CO;) = 1567, v,(CO;) = 1363 and d(CH) = 1375 cm-'1. During heating under evacuation, the intensity of these bands increased.The same features appeared when the sample, maintained at room temperature, was kept in a CO atmosphere for a long time. (c) Two strong complex absorptions near 1190 and 1050 cm-l; their intensity slowly increased with time and with pressure of CO. The following experiments allow determination of the nature of the species leading to the 2100 cm-l band and to the complex absorption in the 1200-1000 cm-l range. (i) On heating at 403 K under vacuum, the band at 2100 cm-l completely disappeared. In the 1200-1000 cm-l range, the intensity of the band at 1190 cm-l was strongly reduced; after heating at 523 K only two bands (1 145 and 1080 cm-l) remained (fig. 2). (ii) Addition of 0, at 298 or 240 K to irreversible species given by CO immediately made the disc white.The complex absorption due to irreversible species between 1200 and 1000 cm-l disappeared, while the band near 2100 cm-l was not affected. Addition of successive small doses of 0, established that the bands at 1190 and 1057 cm-1 were the most sensitive to 0,. In the 1800-700 cm-l range we noted the appearance of bands due to carbonate species [ 1725, 1530, 1478, 1290 and 860 cm-l, see ref. (6)] with a supplementary band at 1750 cm-l. (iii) Addition of successive doses of H,O to irreversible CO species also whitened the disc and decreased the intensity of the 2100 and 120&1000 cm-' bands, which completely disappeared, while bands due to formate species and water adsorption [3740, 3670 (sharp), 3520 (broad) and 670cm-lI appeared.We also noted the appearance of a further unknown species characterized by two bands at 980 and 830 cm-l. (iv) Irreversible preadsorption of CH,OCH, on Th0,-973 did not prevent the appearance of the 2100 and 1200-1000 cm-l bands but the band at 2171 cm-l due to reversible species did not appear. From these experiments we deduced the formation of four types of irreversibly21 8 H, AND CO ADSORPTION ON Tho, I I I I I 1600 1400 1200 1000 800 wavenurn berlcrn-' Fig. 3. 1.r. spectra of species given by successive additions of H, and CO to Tho, : (a) spectra given by H, and outgassing at P, = 850 N rn-, and (b) after addition of 25 pmol g-l CO (the background has been subtracted). absorbed species, in addition to the carbonates and formates; they are characterized by bands at: species I, 2100 cm-l (2044 cm-l); species 11, 1190 and 1055 cm-l (1 158 and 1030 cm-l); species 111, 1165 and 1055 cm-l (1 135 and 1030 cm-l) and species IV, 1145 and 1080 cm-l(l120 and 1055 cm-l).Their thermal stability increased from species I to IV. Complementary experiments were carried out using 13CO; the same features were observed at the wavenumbers reported above in parentheses. The total amount of irreversible species formed is important; after contact with 26 kN rn-, of CO for 0.6 h and outgassing at room temperature for 1 h (Pe = 1.5 x N rn-,), ca. 150 pmol g-l was irreversibly adsorbed. EFFECT OF co ON ADSORBED H, SPECIES Addition of CO to Th0,-973 on which H, has been introduced (P, = 850 N m-2) did not lead to irreversible CO species (fig.3). On the other hand, introduction of CO onto the same sample pretreated with H, and then evacuated at room temperature displaced bands due to H, chemisorption (1 1 15,9 10 and 860 cm-l), while bands due to irreversible CO species (2100 and 1200-1000 cm-l) were formed. In both cases we noted the formation of other species, principally formate species and others characterized by bands at 1235 and 11 10 cm-l. They are not assigned in this paper, but some of them were found on adsorption of HCHO.g EFFECT OF H, ON ADSORBED CO SPECIES On Th0,-973, to which CO had been introduced at room temperature and then evacuated, we observed that adsorption of H, was not prevented. Moreover, the introduction of H, did not perturb the bands due to irreversible CO species (2100 and 1200-1000 cm-l).On heating (423 and 523 K), the intensity of these bands behaved in a similar manner as observed without hydrogen: we deduce that irreversible CO species are insensitive to H,. When the same experiment was carried out without evacuation of CO, we observed a reaction between H, and reversible CO species or CO gas leading to formate and other species characterized by bands at 1265, 1240, 1175, 11 10, 1075 and 922 cm-l.J. LAMOTTE, J-C. LAVALLEY, V. LORENZELLI AND E. FREUND 219 DISCUSSION On alkaline-earth-metal oxides it has been shown that H, dissociates, pro- ducing hydride and hydroxyl groups, which are partially irreversibly adsorbed on eva~uation.~ In particular, very broad bands below 1350 cm-l correspond to hydride groups. Our results on thoria are in agreement, principally from the point of view of the quantity of species formed.Moreover, the 11 15 cm-l band found in the case of thoria correlates well with the polarizing power of the elements (Mg2+, Ca2+, Sr2+ and Th4+). From studies on organometallic compounds, we expect the free Th-H wavenumber to be near 1355 cm-l> lo while a band was found at 11 14 cm-l in bridged compounds.11 They shift to 979 and 802 cm-1, respectively, in deuterated compounds. These results suggest that the species formed on thoria are bridged (Th-H-Th). In the case of heterolytic dissociation, we expect a corresponding v(0H) band. Morterra found it at 3800-3000 cm-l (broad) on Mg0.7 We were not able to find such a broad band. We consider that the two bands observed at 3660 and 3640 cm-l, less sensitive to outgassing than the Th-H-Th species, could arise from another mechanism of adsorption (homolytic dissociation).The adsorption of CO on thoria leads to a orange-yellow colour, as in the case of CO adsorption on highly activated alkaline-earth-metal oxides.8 Preadsorption experiments using CH30CH3 show that the band at 2171 cm-l characterizes reversible species coordinated to Lewis-acid sites. The nature of the species corresponding to the band at 2100 cm-l is difficult to establish. On MgO, Zecchina and coworkers8 associated bands between 2108 and 2064 cm-l with those between 1318 and 1392 cm-l as they were sensitive to oxygen. Work in progress on the adsorption of CO on TiO, tends to show that a band at 21 10 cm-l, resistant to evacuation, corresponds to reduced Ti3+ sites.Such an assignment on Tho, would mean that under our activation conditions a very small number of reduced sites are formed. Another assignment proposed for anatase12 considers the 21 15 cm-1 band to be due to reactive intermediates leading to CO, (carbonate species). Such carbonate species (bands at ca. 1530 and 1330 cm-l) are always found in our experiments. Each species 11-IV is characterized by only two bands: they are situated in the 120Ck1000 cm-l range and are relatively symmetric with an average value close to 1120 cm-l. Note that bands due to carbonate species are also situated symmetrically on either side of an average value (1415-1455 cm-l). We deduce that the structures of species 11-IV are quite similar.The closer their wavenumbers, the higher their thermal stability. Note that the oxygen-sensitive species formed on MgO, CaO or SrO present a much more complex spectrum. We propose two structures, assuming that the bands between 1200 and 1000 cm-l characterize v(C-0) single bonds. As two bands are found for each species, we propose structures with two C-0 groups. (i) The CO molecule may acquire an electron, leading to the radical anion COO-, which can dimerize, giving -0C-CO- species.13 This could be the first member of the oxocarbon anions previously suggested by Zecchina and coworkers.* The two bands characterizing species 11-IV can be assigned to such v(C-0) groups while, from reasons of symmetry, the v(CrC) cannot be observed by i.r. spectroscopy.Such an assignment would agree with the known property of thoria to give higher alcohols (C-C bonds) from the CO + H, intera~ti0n.l~ Moreover, using organo- metallic compounds, Markslo observed that CO reacts with Th-H bonds leading to Th-0-(H)C=C(H)-0-Th groups which may be formed by a similar process. (ii) The other possibility is to consider a negatively charged CO, species. It cannot be a220 H, AND co ADSORPTION ON Tho, carboxylate radical-anion CO;. Such a species has been found by e.p.r. ~pectroscopy,~ but with infrared spectroscopy we expect v(C0) bands at wavenumbers > 1300 cm-l.15 Evidence has recently been obtained16 for CO, ligands in complex molecules showing only one absorption band in the 1200 cm-l range, whose structure has been identified by X-ray diffraction as /"-" M-C \ 0-M. In such a structure the CO, molecule can be formally described as a 'carboxylate' species COi-.Recently, compounds such as Li,CO, have been prepared and characterized by i.r. The splitting between the two v(C0) bands is quite high (440 cm-l) but their mean value is near 1200 cm-l. Moreover, it appears that the 12C + 13C isotopic shifts are 35 and 23 cm-l for the fundamentals v(C0) of Li,CO, situated at 1428 and 987 cm-l, respectively. The present study shows that similar shifts are observed in the case of species 11-IV. Furthermore, note that the anion COi- and the CF, molecule are isoelectronic and that the spectrum of the latter has bands at 1222 and 1102 cm-l.lS Other molecules or ions such as NO;, 0, and ClOi also have bands in the 1300-1000cm-1 range.18 The formation of such COi- species would probably involve 0,- basic sites: [<I (CO) Th4+02- - Th4+ C It is difficult to choose between these two interpretations; the colour was explained by the former, as Zecchina and coworkers invoke the formation of negative CO polymer species on alkaline-earth-metal oxides.They could result from a dispropor- tionation reaction : (n+;)Co+xo2-'-co:f(CO):-. x 2 On thoria such a reaction does not seem very important as the bands due to carbonate species were very weak when CO was adsorbed at room temperature (fig. 2). The latter interpretation is more simple although the formation of COi- species has not been reported on metallic oxides. Note, however, that Tho, is a strongly basic oxide and the surface oxide ions may act as strong electron donors to adsorbed molecules.From our study of the effect of the preadsorption of one molecule (H, or CO) on the other we deduce that different sites are responsible for the dissociation of H, molecules and for the formation of irreversible CO species (species 11-IV): the latter do not prevent dissociation of H, and are not reduced by H,; on the other hand, it appears that Th-H-Th species react with reversible CO species or CO gas, leading to species similar to those formed by the adsorption of HCHO. They could be considered as intermediates in the synthesis of methan01.~ ' E. Druet, Ph.D. Thesis (I.F.P., Rueil Malmaison, 1982). * R. Bardet, J. Thivolle-Cazat and Y. Trambouze, C.R. Acad. Sci., 1981, 292, 883.J. C. Lavalley, J. Saussey and T. Rais, J . Mof. Cataf., 1982, 17, 289. P. Pichat, J. Veron, B. Claude1 and M. V. Mathieu, J . Chim. Phys., 1966, 63, 1026. W. S. Brey, R. B. Gammage and Y . P. Virmani, J . Phys. Chem., 1971, 75, 895.J. LAMOTTE, J-C. LAVALLEY, V. LORENZELLI AND E. FREUND 22 1 J. Lamotte, J. C. Lavalley, E. Druet and E. Freund, J. Chem. Soc., Faraday Trans. I, 1983, 79, 2219. ' S. Coluccia, F. Boccuzzi, G. Ghiotti and C. Morterra, J. Chem. SOC., Faraday Trans. 1, 1982,78,2111. E. Guglielminotti, S. Coluccia, E. Garrone, L. Cerruti and A. Zecchina, J. Chem. SOC., Faraday Trans. I , 1979, 75, 96; S. Coluccia, E. Garrone, E. Guglielminotti and A. Zecchina, J. Chem. SOC., Faraday Trans. I , 1981, 77, 1063. To be published. lo P. J. Fagan, K. G. Moloy and T. J. Marks, J. Am. Chem. Soc., 1981, 103, 6959. l 1 J. M. Manriquez, P. J. Fagan and T. J. Marks, J. Am. Chem. SOC., 1978, 100, 3939. l 2 K. Tanaka and J. M. White, J. Phys. Chem., 1982, 86, 4708. l 3 P. W. Lednor and P. C. Versloot, J. Chem. SOC., Chem. Commun., 1983, 284. l4 E. M. Cohn, in Catalysis, ed. P. H. Emmett (Reinhold, New York, 2nd edn, 1961), vol. IV. l5 G. Busca and V. Lorenzelli, Mater. Chem., 1982, 7, 89. l6 C. R. Eady, J. J. Guy,B. F. G. Johnson, J. Lewis, M. C. MalatestaandG. M. Sheldrick, J. ChemSoc., l7 Z. H. Kafafi, R. H. Hauge, W. E. Billups and J. L. Margrave, J. Am. Chem. SOC., 1983, 105, 3886. " K. Nakamoto, Infrared and Raman Spectra of Inorganic and Coordination Compounds (Wiley, New Chem. Commun., 1976, 602. York, 3rd edn, 1978). (PAPER 4 / 7 4 )

ISSN:0300-9599    

DOI:10.1039/F19858100215

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1985, 81, 223-232 Enthalpy and Volume Changes on Mixing Diethylene Glycol Di-n-alkyl Ethers with Diethylene Glycol Dimethyl Ether or n-Alkanes BY JAWAD K. H. AL-KAFAJI,~ ZAITUN ARIFFIN, JEREMY COPE AND COLIN BOOTH* Department of Chemistry, University of Manchester, Manchester M 13 9PL Received 10th May, 1984 Enthalpy and volume changes on mixing diethylene glycol di-n-alkyl ethers1 (alkyl- oxyethylene-alkyl triblock oligomers) with either diethylene glycol dimethyl ether [oligo(oxy- ethylene)] or the C,,, C,, and C,, n-alkanes have been measured at 3 1 "C. The results for these pseudo-binary mixtures are compared with those for binary systems, i.e. oligo- (oxyethy1ene)s + n-alkanes. Oligoethylene glycol alkyl ethers have a wide application as non-ionic surfactants and are also of interest because microphase separation, in both concentrated aqueous solution and the pure state, leads to a variety of interesting structures.l We have investigated2 the structural chemistry of a number of homogeneous oligoethylene glycol di-n-alkyl ethers and have shown that their component blocks are separated in the solid state but mixed in the molten state.Consequently a contribution to the enthalpy of fusion of the solid comes from the enthalpy of mixing of the component blocks in the melt, which means that the thermodynamics of fusion of these systems can be understood in detail only if the thermodynamics of mixing in oligo(oxyethy1- ene) + oligo(methy1ene) systems is first investigated. It is known3 that the heat of mixing per unit volume in the oligo(oxyethy1ene) + n- alkane system is large and positive: e.g.ca. 12 J cmP3 for an equi-volume mixture. Consequently the diminution of the entropy of mixing per unit volume as chain length is increased results in immiscibility at low chain lengths3 and mixing experiments have to be carried out with the lower oligomers. (In the presence of water, even in small amounts, this incompatibility is enhanced.) The thermodynamics of mixing will be changed when the oxyethylene and methylene components are blocks within the same chain rather than separate chains, and it is probable that the enthalpy of mixing is affected, albeit in a less obvious manner than the entropy of mixing. We have sought experimental information on this point by investigating the mixing of diethylene glycol dimethyl ether or n-alkanes with diethylene glycol di-n-alkyl ethers, H(CH,),(OCH,CH,),O(CH,),H, where n = 4 or 6.EXPERIMENTAL NOTATION The n-alkanes are denoted C,, (dodecane), C,, (tetradecane) and C,, (hexadecane). The diethylene glycol diethers are denoted n-2-n, where n is the alkyl chain length and 2 is the nominal oxyethylene chain length in oxyethylene units. Thus diethylene glycol di-n-butyl ether is denoted 4-2-4. t Present address : Chemistry Department, College of Education, University of Baghdad, Baghdad, Iraq. $ Systematic name: a-(n-alky1)-o-(n-a1koxy)-di(oxyethy1ene). 223224 ENTHALPY AND VOLUME CHANGES ON MIXING MATERIALS Several samples were obtained from commercial sources: 1-2- 1 from Hopkin and Williams Ltd; 4-2-4, C,, and C,, from Fluka A.G.; C,, from B.D.H.Ltd. Sample 6-2-6 was prepared from diethylene glycol (Shell Chemical Co. Ltd) and 1 -bromohexane by Cooper's modification of the Williamson ether synthesis, as described elsewhere.2* * The yield was ca. 20% in each of two preparations: the reason for the low yield is not understood. Characterisation of the prepared samples, by i.r. and lH n.m.r. spectroscopy, showed conversions of 98% or better. When necessary the liquids were fractionally distilled to obtain samples of acceptable purity ( > 99 wt % , as determined by gas-liquid chromatography). METHODS The experimental methods used for determination of specific volume (pyknometry), volume change on mixing (dilatometry) and enthalpy change on mixing (calorimetry) have been de~cribed.~ All mixing experiments were carried out at 31 "C.RESULTS SPECIFIC VOLUMES Specific volumes (u,,) at 31 "C are listed below. Those for the diethylene glycol diethers were determined in this work to kO.002 cm3 g-l; those for the n-alkanes are taken from the literat~re:~ sample 1-2-1 4-2-4 6-2-6 C,, C14 C,, uSp/cm3 g-l 1.072 1.144 1.156 1.350 1.324 1.306. ENTHALPY AND VOLUME CHANGES Enthalpy changes at 31 "C on mixing the block oligomers with either diethylene glycol dimethyl ether or n-alkane were measured for the following systems: 4-2-4+ 1-2-1, C12, C14 or C,, and 6-2-6+ 1-2-1. Volume changes at 31 "C were measured for the system 4-2-4 + 1-2-1 or C12. The results obtained are listed in table 1. MIXING AND DILUTION The enthalpy of mixing n, moles of segments of oligo(oxyethy1ene) (component 1) with n-alkane (component 2) so as to form a mixture of volume fraction 4, may be writ ten When a homo-oligomer is added to a block oligomer the process is analogous to the dilution of a mixture by one of its components.When a mixture of composition specified by volume fraction 4; is diluted with An, moles of segments of homo-oligomer 1 to form a mixture of volume fraction 4;, the enthalpy change may be written AH, = RTn,42x. (1) A similar expression may be written for the enthalpy change on dilution of the mixture by An2 moles of homo-oligomer 2, i.e. AHd = RTAn, 4; 4: c,. (3) The enthalpy parameters x, el and c2 are equal only if they are independent of concentration. If the mixing parameter x has the form (4)J.K . H. AL-KAFAJI, Z. ARIFFIN, J. COPE AND C. BOOTH 225 Table 1. Enthalpy and volume changes on dilution of block oligomers with homo-oligomers ~(4-2-4)lg 0.092 0.181 0.268 0.359 0.444 0.525 0.634 0.712 0.794 ~ ( 1 - 2 - l ) l g 0.866 0.775 0.683 0.584 0.489 0.395 0.301 0.215 0.111 AH/J 0.8, 1.4, 1.9, 2.2, 2.3, 2.2, 2.0, 1.6, 0.9, w(C,,)/g 0.102 0.199 0.276 0.338 0.431 0.490 0.579 0.657 0.733 AHIJ 1.5, 2.5, 3.2, 3.6, 3.8, 3.6, 3.2, 2.6, 1.8, ~(4-2-4)lg 0.884 0.864 0.729 0.594 0.525 0.402 0.313 0.217 0.125 ~(4-2-4)lg 0.826 0.737 0.646 0.550 0.472 0.381 0.295 0.204 0.155 w(C,,)/g 0.097 0.175 0.251 0.323 0.402 0.476 0.554 0.631 0.706 AH/J 1.5, 2.4, 3.1, 3.4, 3.7, 3.6, 3.3, 2.7, 1.7, w(C,,)/g 0.098 0.178 0.253 0.332 0.412 0.479 0.561 0.637 0.714 ~(4-2-4)lg 0.824 0.737 0.650 0.561 0.473 0.380 0.296 0.205 0.116 AHIJ 1.5i 2.2, 3.2, 3.6, 3.8, 3.6, 3.4, 2.9, 1.9, ~(6-2-6)/g 0.106 W( 1-2- l)/g 0.850 AH/J 1.5, ~(4-2-4)lg 0.90 ~ ( 1 - 2 - l ) l g 5.01 .4V/mm3 4., w(C,,)/g 1.32 AV/mm3 13., ~(4-2-4)/g 4.20 0.291 0.463 0.633 0.809 0.645 0.468 0.288 0.104 3.1, 3.7, 3.2, 1.5, 3.09 3.24 3.94 4.86 5.02 4.75 3.25 1.01 4.92 4.47 1.91 1.12 2.24 2.06 3.91 3.32 19., 18., 22.:, 16., 1 l., lo., 9., 4., then, since ( 5 ) A similar development, more appropriate to the oligo(oxyethy1ene) + n-alkane ~ y s t e m , ~ starting from x = X a + X b 4 2 + X c 4 ; (8) (9) leads to 5, = ka -Xb 4; -Xc 4; 4;) + k b - X c 4;) 4; +XC(~;)' Values of x and calculated for first- and second-degree dependences of x on 42, with an initial concentration for dilution of 4; = 0.5, are illustrated in fig.1. A discontinuity in [ at 4; is a feature of any system in which x varies as illustrated. * AHd = AH; - AH; = R Q a ( n ; 4; - n; 4;) +x,,(n; 4i2 - n ; 42) = RTCy,[n;(& - 4;) + Anl 43 +x,[n;(4z2 - @ ) + A n , 4 3 ) which, with some simplification, gives eqn (6). A similar development, from x = ka+xb)+xb4,, gives eqn (7).226 CH,CH,CH, 2 1 .o I CH2CH2CH3 2 CH,OCH,CH,OCH,CH,OCH, 1 0. 0 . 0 ENTHALPY AND VOLUME CHANGES ON MIXING 0.5 4 2 1.0 0 0 -5 4 2 1 .o Fig. 1. Comparison of mixing k) and dilution ([) parameters for model systems. (a) x = 0.7+0.24,; [, and [, given by eqn (6) and (7) with 4; = 0.5. (b) x = 0.7-0.2~$,+0.4#;; and [, given by eqn (9) and (10) with 4; = 0.5. TREATMENT OF RESULTS In order to use the equations developed above the composition of the block oligomer must be defined.For block oligomers similar to those used here, the well established correlation of chemical shift in n.m.r. and the crystalline structures2 and Raman LAM frequencies,6 in addition to chemical logic, strongly support the disposition of the CH, group adjacent to the ether oxygen to the oxyethylene block. Hence, for example, we assign sample 4-2-4 the composition specified by a weight fraction w1 = 0.605. This procedure is consistent with the use of diethylene glycol dimethyl ether as the homo-oligomer. However, the block oligomer composition must be defined by volume fraction forJ. K. H. AL-KAFAJI, Z. ARIFFIN, J. COPE AND C. BOOTH 227 1 .o 0.8 0 - 6 0 0.5 1.0 @; Fig.2. Enthalpy parameters for dilution of block oligomer 4-2-4 by (0) oligo(oxyethy1ene) 1-2-1 (Cl) or (a) dodecane (Cz). The points are calculated from the enthalpy results (see table 1) using eqn (2) and (3) with the specific volume ratio R, (i.e. vl, s p / ~ 2 , sp) equal to (0, A) 0.70 or (e, A) 0.76. The curves are calculated from x [eqn (14)] using eqn (9) and (10) with R, = 0.70 (dashed curve) or 0.76 (full curve). any theoretical development in which core volume is Consequently it is necessary to estimate the ratio of specific volumes, R, = ul,sp/v2,sp. (a) A simple assumption is that the oxyethylene and alkyl blocks have specific volumes in the same ratio as those of their homo-oligomers of the same overall molecular length (i.e. both 15 chain atoms for sample 4-2-4).(b) An alternative assumption is that the oxyethylene and alkyl blocks have specific volumes in the same ratio as those oftheir homo-oligomers of the same block length [i.e. 9 chain atoms for oligo(oxyethylene), 6 chain atoms for n-alkane, for sample 4-2-41. Since the specific volumes of the homo-oligomers are linear functions of reciprocal chain length, interpolation of published3v results is straightforward. For sample 4-2-4 the ratio R, = ul, sp/vz, sp is found to be ((I) 0.76 and (b) 0.70, which leads to values of 4; of (a) 0.538 and (b) 0.517. It is likely that these two values of R, represent its upper and lower extremes. The consequences of error in this assumption are illustrated in fig. 2. It is seen that the values of cl and c2 are changed between assumptions (a) and (b), but that the effect is small, ca. 5% for [.The complete set of values of c is shown in fig. 3. Experimental errors in the determination of AHare estimated to be kO.3-+ 0.5 J; the consequential errors in c are Volume changes may be treated in an equivalent manner to enthalpy changes. For mixtures of homo-oligomers 0.02- 0.05. The segment volume is 39.4 cm3 mol-1 AVm = V,n14,A (1 1)228 1 . c 5 0.e 0.6 0 ENTHALPY AND VOLUME CHANGES ON MIXING / / 0.5 4; 1.0 Fig. 3. Enthalpy parameters for dilution of block oligomers by homo-oligomers. The specific volume ratio R, (i.e. vl,sp/uz,sp) = 0.76 for both block oligomers. The points are calculated from enthalpy results (see table 1) using eqn (2) and (3) for block oligomer 4-2-4 plus (0) 1-2-1, (A) C12, (A) C,, or (V) C,, and for block oligomer 6-2-6 plus (0) 1-2-1.The curves are calculated from x [eqn (14)] using eqn (9) and (10) for 4-2-4 (full curves) and 6-2-6 (dashed curve). where & is the segment volume; for dilution of block oligomer by oligo-oxyethylene and for dilution of block oligomer by n-alkane The same assumptions with respect to composition and specific volume as used earlier (i.e. R, = 0.76, w, = 0.605) lead to the values of r, and 5, shown in fig. 4. Experimental errors in the determination of AVd are estimated to be ca. & 1 mm3; the consequential errors in < are f0.003-+0.005. COMPARISON OF BLOCK- AND HOMO-OLIGOMER SYSTEMS Values of Cl and c2 calculated using eqn (9) and (1 0) with the x parameter established for diethylene glycol dimethyl ether (sample 1-2-1) mixed with the n-alkanes C,, to C17, i.e.(14) are plotted against 4; in fig. 2 and comparison made with the observed values of Cl and C2 for the block oligomer system 4-2-4+ 1-2-1 or C12. It can be seen that the x = 0.797 - 0.3444; +0.525(4g)2J. K. H. AL-KAFAJI, Z. ARIFFIN, J. COPE AND C. BOOTH 229 0.06 t 0.04 0.02 0 0 . 5 1 .o @; Fig. 4. Volume parameters for dilution of block oligomer 4-2-4 by (0) oligo(oxyethy1ene) 1-2-1 (CJ or (A) dodecane (c2). The points are calculated from the volume results (see table 1) using eqn (12) and (13) with the specific volume ratio R, (ie. = 0.76. The curves are calculated from ,I [see eqn (5) of ref. (3)] using equations analogous to eqn (9) and (10). calculated values of c are insensitive to choice of R, and that the discrepancies between observed and calculated values of [ are sufficiently large that the uncertainty introduced by the need to specify R, (see fig.2) is unimportant. A more complete comparison between the block-oligomer and the homo-oligomer systems is made in fig. 3 and 4. Only the assumption (a) for R, is illustrated: the results of comparison under assumption (6) do not differ significantly. The comparison between block-oligomer and homo-oligomer systems can be made in another way. The values of the parameters el and c2 can be fitted by quadratics in 4; (by least squares) and, from the resulting coefficients, values of parameter x can be calculated using rearranged eqn (9) and (10). A comparison of x, obtained in this way for the system 4-2-4+ 1-2-1 or CI2, with x obtained directly from mixing the homo-oligomers 1-2-1 and C,, [ i e .eqn (14)] is shown in fig. 5. This calculation of x is unstable for scattered results and the plot may not be correct in detail. The reverse calculation (of or [, from x) is soundly based as the equation for x is based on many data point^.^ DISCUSSION In the mixing of oligo(oxyethy1ene)s and alkanes, the variations of enthalpy and volume parameters and 1) with concentration (42) are similar in nature. It has been ~ h o w n , ~ by use of Flory’s theory,8 that this is because equation-of-state effects make230 ENTHALPY AND VOLUME CHANGES ON MIXING 1 .o 0 . 8 X 0.6 0 0 . 5 41 1 .o Fig. 5. Enthalpy parameters for the oligo(oxyethy1ene) plus oligo(methy1ene) system. The full curve is x from the results of mixing the oligo(oxyethy1ene) 1-2-1 with the n-alkanes C,, to C,, [eqn (14)].The dashed curve is x calculated from the results of diluting the block oligomer 4-2-4 with 1-2-1 or C,, (i.e. Cl or rz, specific volume ratio R, = 0.76; see fig. 3). only a small contribution to the mixing parameters (< 20% of A and < 1 % of x) so that the volume change is determined principally by the energy change. Consequently the similar variations of enthalpy and volume parameters found in this work (cf. fig. 3 and 4) are as expected. In the following discussion we concentrate on the enthalpy parameters, for which equation-of-state effects are very small. The enthalpy parameters found on mixing the block oligomers with homo-oligomers are lower in value than those found on mixing two homo-oligomers: i.e.with regard to enthalpy, the block oligomers are more compatible with each of the homo-oligomers than the homo-oligomers are with each other. Moreover, the concentration dependence of the enthalpy parameters is less for the block-oligomer + homo-oligomer systems than for the homo-oligomer + homo-oligomer systems. It has been noted earlier3 that a value of the segment surface-area ratio (Slz = SJS,) near 0.8 serves to represent the upward trend in x with dZ, when it is assumed8 that the interaction energies are additive in segment surface area. The curvature in ~ ( 4 ~ ) cannot be described by a theory8 based on random mixing of chains and has been ascribed3 to quasi-chemical interaction of segments caused by the large positive interchange energy.A simple model of the molecules as cylindrical rods with dimensions similar to those established by crystallography* leads to a lower average surface area per segment for the central oxyethylene block compared with a homo-oligo(oxyethylene), because of the continuity of the chain (i.e. the missing ends), and to an unchanged average surface area per segment for the end n-alkyl blocks. * Chain length per chain atom: oxyethylene, 0.0928 nm; methylene 0.1 27 nm. End contribution ca. 1 nm. Cross-sectional area: oxyethylene, 0.214 nm2; methylene, 0.184 nm2.J. K. H. AL-KAFAJI, Z. ARIFFIN, J. COPE AND C. BOOTH 23 1 A discrepancy of 10-20% is likely. As a consequence, compared with the homo- oligomer + homo-oligomer systems, the overall surface area per segment in the block-oligomer + homo-oligomer systems will be lower and the surface-area ratio S,, will be higher.Thus the predicted effect on x of the changed molecular structure is a lower value for the block-oligomer system and a reduced concentration dependence. As to the curvature of x if, as has been ~uggested,~ this is due to the quasi-chemical interaction of segments, then linking the unlike blocks within the same chain will promote random mixing and reduce the effect. EFFECT OF n-ALKANE CHAIN LENGTH For mixtures of diethylene glycol dimethyl ether (1 -2- 1) with n-alkanes, values of x show3 no differences outside experimental error as the n-alkane is varied between C,, and C17.The present results show an effect caused by the chain length of the n-alkane (see fig. 3). The enthalpy change on mixing equal volumes of the two homo-oligomers is3 ca. 12 J cmP3, whereas that found on diluting the block oligomer 4-2-4 with an equal volume of n-alkane is < 4 J ~ m - ~ (see table 1). Consequently small effects caused by chain length may be more apparent in the present results than in the earlier ones. It has been shownlo9 l1 that correlation of short-range molecular order (c.m.0.) in oligomers can be changed on mixing and that this can lead to a positive contribution to the enthalpy of mixing. The topic has been reviewed recently.,, The effect has been formulatedlO as an extension of Flory's theory by writing the interchange energy parameter thus incorporating contributions from interchange of contacts ( X g ) and from changes in correlation of molecular order [X,,(c.m.o.)].Redefinition of A',, in this way does not, in itself, change the predicted variation of x with 4, in a binary mixture but it may effect ~ ( 4 , ) in a pseudo-binary mixture of the type under investigation here. Parameter X,,(c.m.o.), and consequently x, will increase with increasing chain length of oligomerloy l1 whenever correlation of molecular order is lessened on mixing. The results shown in fig. 3 for sample 4-2-4 diluted with the n-alkanes C,,, C,, and C,, are consistent with this. The change in on changing from C,, to C,, as diluent is ca. +0.05. This corresponds to a value of X12(c.m.o.) z 5 J cm3, which is a reasonable value.l0-l2 However, correlation of molecular order does not provide an explanation for the form of ~ ( 4 , ) which, for the oligo(oxyethy1ene) + n-alkane system, comprises3 a symmetrical deviation from theory8 proportional to the product of volume fractions 4, 42.X,, = X g +X,,(c.m.o.> (1 1) We thank Dr R. H. Mobbs for guidance in the preparative work and Dr R. 0. Colclough for helpful discussions. G. T. J. Tiddy, Phys. Lett., Phys. Rep., 1980, 57, 1. R. C. Domszy and C. Booth, Makromol. Chem., 1982, 183, 1051; H. H. Teo, T. G. E. Swales, R. C. Domszy, F. Heatley and C. Booth, Makromol. Chem., 1983, 184, 861. J. K. H. Al-Kafaji and C . Booth, J . Chem. SOC., Faraday Trans. 1, 1983, 79, 2695. D. R. Cooper and C. Booth, Polymer, 1977, 18, 164. R. A. Orwoll and P. J. Flory, J. Am. Chem. Soc., 1967, 89, 6814. T. G. E. Swales, H. H. Teo, R. C. Domszy, K. Viras, T. A. King and C. Booth, J . Polym. Sci., Polym. Phys. Ed., 1983, 21, 1501. ' I. Prigogine, A. Bellemans and V. Mathot, The Moleculur Theory of Solutions (North Holland, Amsterdam, 1957).232 ENTHALPY AND VOLUME CHANGES ON MIXING * P. J. Flory, Discuss. Faraday SOC., 1970, 49, 7. D. Patterson and G. Delmas, Discuss. Faraday SOC., 1970, 49, 98. lo V. T. Lam, P. Picker, D. Patterson and P. Tancrede, J. Chem. Soc., Faraday Trans. 2, 1974,70, 1465; M. T. Croucher and D. Patterson, J . Chem. SOC., Faraday Trans. 2, 1974, 70, 1479; P. Tancrede, D. Patterson and V. T. Lam, J . Chem. SOC., Faraday Trans. 2, 1975, 71, 985. P. Tancrede, P. Bothorel, P. de St. Romaine and D. Patterson, J . Chem. SOC., Faraday Trans. 2, 1977, 73, 15; P. Tancrede and D. Patterson, J. Chem. SOC., Faraday Trans. 2, 1977, 73, 29. l 2 A. Heintz and R. N. Lichtenthaler, Angew. Chem., Znt. Ed. Engl., 1982, 21, 184. (PAPER 4/757)

ISSN:0300-9599    

DOI:10.1039/F19858100223

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. 1, 1985,81, 233-239 Reactions of Iron@) Protoporphyrin with Strongly Reducing Free Radicals in Aqueous Solutions A Pulse-radiolytic Study BY YACOV SOREK, HAIM COHEN* AND DAN MEYERSTEIN* Nuclear Research Centre, Negev, Israel and Department of Chemistry, Ben Gurion University of the Negev, Beer-Sheva 84105, Israel Received 1 1 th May, 1984 Iron(I1) protoporphyrin. Fe'IPP, is reduced by eLq and Cog- in aqueous solutions to Fe'PP, which might be present in aqueous solutions as FeII'PP-H. The spectrum of the short-lived intermediate and the kinetics of its decomposition into FelIIPP + H, are reported. The free radicals 'CH,OH, 'CH(CH,)OH and 'C(CH,),OH react with FeIIPP forming Fe'l'PP-CR,R,OH ; the spectra of these intermediates, which decompose into F;elIIPP+ HCR,R,OH, are reported.Redox reactions of iron porphyrin complexes have been the subject of extensive studies because of their biological Within this framework, the reactions of iron(I1) porphyrins with several aliphatic free radicals in aqueous solutions have recently been ~tudied.~ The results indicated that for oxidizing aliphatic radicals, e.g. 'CH, and 'CHCICF,, the products of the reaction are a-bonded alkyl iron complexes :9c There is little information available on the reactions of ferrous porphyrins with reducing radicals, e.g. a-hydroxyalkyl free radicals, CR,R,OH. The reaction with 'C(CH,),OH was claimed to give unidentified Furthermore, it has been suggested that the product of the reaction of e& with iron(I1) porphyrins is the iron(I1) porphyrin anion-radical complex, as ferrous ions are difficult to reduce.1° In unpublished workll it was suggested that (CH,),CO'- radicals reduce the iron(I1) porphyrins to an iron(1) complex or to the anion-radical species, which are then back oxidized to the iron(i1) form.On the other hand electrochemical reduction of iron(I1) porphyrins in DMF yields first the iron(r) porphyrins and iron(r) porphyrin anion radicals are formed ca. 0.5 V more cathodically.8 We have decided to study the reactions of the strongly reducing free radicals eFq, Cog-, 'C(CH,),OH, 'CH(OH)CH, and *CH,OH with iron(1r) protoporphyrin, FeI'PP. The results indicate that eiq and COH- reduce FellPP. The reduced intermediate reacts with water yielding FeIIIPP + H,. The a-hydroxy free radicals oxidize Fe'IPP to FelIIPP via an intermediate which is tentatively identified as FeT1IPP-CR, R,OH.FeI1P+ 'CR,R,R, -+ FelllP-CR,R,R,. (1) EXPERIMENTAL MATERIALS Solutions of Fe"PP were obtained by reducing chlorohemin, FeII'PP, from the Nutritional Biochemical Corporation. Stock solutions of Fe'I'PP at pH 11.0 were reduced by the introduction of Adams catalyst, PtO,, and bubbling dihydrogen for ca. 6 h. The stock Fe"PP 233234 REDUCTION OF IRON(II) PROTOPORPHYRIN solution thus obtained was transferred into an all-glass syringe through a glass sinter. The completion of the reduction was checked spectrophotometrically. The spectrum of Fe"PP thus obtained is identical to that reported in the literature.' Thus one can reduce FeI**PP by this technique, which is preferable as it does not introduce any additives to the solution.Solutions with the required composition were prepared by mixing this stock solution with Ar- or N,O-saturated solutions containing all the other chemicals required, using the syringe technique. All other chemicals were of AnalaR grade and were used without further treatment. All solutions were prepared with distilled water which was further purified using a Millipore filtering system. IRRADIATIONS The pulse-radiolytic experiments were carried out using the electron linear accelerator of the Hebrew University of Jerusalem. 5 MeV, 200 mA, 0.1-0.5 ,us pulses with a dose of 10&500 rad per pulse were applied. A detailed description of the experimental set-up and the techniques used for data analysis has been published previously.'2 Steady-state irradiations were carried out in a 6oCo y-source with a dose rate of 2.5 x lo5 rad h-l.ANALYSIS The yield of dihydrogen was determined by gas chromatography using a thermal-conductivity detector. The yield of H, as a G* value was calibrated by irradiating, with the same dose, N,O-saturated blank solutions containing I x mol dm-3 NaBr for which G(H,) = 0.40.13 PRODUCTION OF THE DESIRED FREE RADICALS The radiolysis of water may be summed up by the equation Y . e- H,O - eiq, OH, H, H,, H,O,, H30+ with the yield of the products being GeG = Go, = 2.65, GH = 0.60, GH2 = 0.45 and GH2O2 = 0.75.13 In solutions containing high concentrations of solutes the free-radical yield is often larger and the yield of H, and H,O, is smaller.The free radicals formed are homogeneously distributed in the solution within < 100 ns of the radiation being absorbed. When studying reactions of eiq one follows the bleaching of the absorption band due to e& in the red region. Usually C(CH,),OH is added to the solution in order to scavenge the strongly oxidizing OH radicals : OH/H + C(CH,),OH -+ 'CH,C(CH,),OH + H,O/H, (3) k,, = 4.2 x lo8 dm3 mo1-l s-' and kH = 8 x lo4 dm3 mol-I s-'.15 The CH,C(CH,),OH radicals are poor reducing agents and are often unreactive towards other solutes and decompose via dimerization, a diffusion-controlled reaction. However, one has to check in each system studied that the 'CH,C(CH,),OH radicals are indeed not reactive. The hydrated electron reacts with N,O according to eiq+N,O -+ N,+O- k, = 8.7 x lo9 dm3 mol s-l l6 (4) followed by O-+H,O -+ OH+OH- k, = 1.7 x lo6 dm3 mol-' s-l. ( 5 ) Thus in neutral or alkaline N,O-saturated solutions, [N,O] = 2.2 x lo-, mol dmP3, > 90% of the primary free radicals are transformed into OH radicals, the rest being hydrogen atoms.When formate ions are added to the solutions the following reactions occur: OH/H + HCO; + COh- + H,O/H, (6) ko, = 3 x lo9 dm3 mol-' s-"* and kH = 1.3 x lo8 dm3 mol-l S-' l5 * G values denote the yield of radiolytic products in units of the number of product molecules per 100 eV absorbed in the sample.Y. SOREK, H. COHEN AND D. MEYERSTEIN 23 5 thus forming the C0,- free radical, which is a powerful reducing agent: E"(CO,/CO,-) = -2.0 V." In solutions containing alcohols the following reactions occur : OH/H + R,R,CHOH -, 'CR,R,OH +H,O/OH ( 7 ) (in this study R, and R, are H or CH,).The specific rates for the reactions with OH are 7 x los, 1.7 x lo9 and 2 x lo9 dm3 mol-l s-l l 8 and with H atoms 1 . 6 ~ lo6, 2.5 x lo7 and 1.8 x los dm3 mol-l s-' l9 for CH,OH, CH,CH,OH and (CH,),CHOH, respectively. [In the ethanol and propan-2-01 solutions ca. 13% of the free radicals are formed by 8-hydrogen abstraction, i.e. the radicals formed are 'CH,CH,OH and 'CH,C(CH,)HOH.] RESULTS AND DISCUSSION REDUCTION OF FexxPP mol dm-3 FeIIPP and 0.1 rnol dm-3 C(CH,),OH are irradiated by a short pulse no spectral changes in the region 450-700 nm (where the solutions are transparent) are observed, suggesting that the reaction of 'CH,C(CH,),OH with FeLIPP is too slow to compete with the dimerization of the free radicals.When identical solutions saturated with Ar are irradiated a short-lived transient absorbing in the 450-700 nm region is formed at the same rate as the bleaching of the absorption of e;,. From the dependence of the observed rate on [FeIIPP] the specific rate of the reaction When N,O-saturated solutions at pH 10.0 containing 2 x FeIIPP + eiq + FeIPP (8) k, = (6 1) x lo9 dm3 mol-l s-l was determined. Fig. 1 presents the spectrum of FeIPP or FeIIIPP-H which might be formed via FeIPP + H20 -+ FeIIIPP-H + OH,, if reaction (8) is the rate-determining step. FeIIIPP-H denotes the FeIII-hydride complex of FeIII-PP by analogy with (H20)5FeH2+.20 (The reasons for suspecting that the observed product might be FeIIIPP-H are discussed below.) The formation of the short-lived transient is followed by a first-order process with a rate of (2.5f0.5) x lo3 s-l, independent of the concentrations of the solute, pulse intensity and wavelength.The product of the process has an absorption spectrum identical to that of Fe'IIPP, suggesting that the reaction occurring is (9) H*O FeIPP or FeIIIPP-H - FeIIIPP + H, (10) k,, = (2.5k0.5) x lo3 S-l. An analogous reaction is not expected if the reduction product is FeIIPP'- and not FeIPP. Indeed when Ar-saturated solutions containing 4 x lo-, mol dm-3 FeIIPP and 0.1 mol dm-3 C(CH,),OH are irradiated by a dose of 2.5 x lo5 rad, dihydrogen is formed with a yield of G(H,) = 2.95. This yield is slightly lower than GIIz + GH + GeBq = 3.7 but considerably higher than G(H2) = 1.05, GHz + GH, observed in the absence of FeIIPP.The smaller than expected yield of dihydrogen might be due to a reaction between H atoms and FeIIPP, which does not yield dihydrogen and competes with reaction (7), or to competition between the reactions of eiq and H with FeIIPP and FeIIIPP, as the latter complex accumulates during the irradiation. Alternatively the reaction of FeIPP with FeIIIPP might decrease the dihydrogen yield.236 REDUCTION OF IRON(II) PROTOPORPHYRIN I 10 5 1 4 50 550 650 i/nm Fig. 1. Spectra of products observed after a pulse producing 6.1 pmol dmP3 of free radicals [Ar-saturated solution at pH 10.1, 2 x loP5 mol dmP3 FeI'PP, 0.1 mol dmP3 (CH,),C-OH, 12.5 cm optical path]: 0, 40 ps after the pulse and A, 2 ms after the pulse.When N,O-saturated solutions containing (2-10) x lop5 mol dm-3 FellPP and 0.1 mol dm-3 HC0,Na at pH 10.0 are irradiated two consecutive reactions are observed. The first obeys a pseudo-first-order rate law, the rate being proportional to [FeIIPP]. The second process obeys a first-order rate law with k = (2.0k0.5) x lo3 s-l, suggesting that the reactions observed are FeIIPP + C0;- + FeIPP + CO, (1 1) k,, = (8 2) x lo7 dm3 mol-1 s-l followed by reaction (10). This is confirmed, as the yield of dihydrogen formed when an N,O-saturated solution containing 4 x lop3 mol dm-3 FeIIPP and 0.1 mol dmP3 HC0,Na at pH 10 is irradiated in the y source by a dose of 2.5 x lo5 rad is G(H,) = 5.75, i.e. it approaches GH1+ GeLq + Go, + G, = 6.35. REACTIONS OF FeIIPP WITH 'CR,R,OH FREE RADICALS When N,O-saturated solutions containing (1-3) x mol dm-3 FeIIPP and 0.1 mol dm-3 alcohol (methanol, ethanol or propan-2-01) at pH 10.0 are irradiated two consecutive processes are observed.* The first obeys a pseudo-first-order rate law, * Identical results are obtained in the pH range 9-1 3 and at pH 13.0 in solutions containing 6.5 mol dm-3 C,H,OH.As under the latter conditions Fe"'PP is in its monomeric form the results seem to be independent of the degree of aggregation of Fe"PP.Y. SOREK, H. COHEN AND D. MEYERSTEIN 237 Table 1. Spectral features of the intermediates observed, their specific rates of formation and decomposition and measured dihydrogen yield solutea saturation AIb/nm ~ , ~ / d m , mol-' cm-' A,b/nm Ezc/dm3 mol-' cm-l kfd/dm3 mol-1 s-' kde/S-' G(H,)r 0.1 mol dmP3 (CH,) ,C HOH NZO 570 11 000 660 5 300 6 x lo8 300 50 0.88 0.1 mol dm-3 CH,CH,OH N2O 560 13 000 680 4 200 4 x lo8 250 & 50 0.88 0.1 mol dm-, 'CH,OH N2O 575 9 000 680 4 800 2.5 x lo8 250 f 50 0.95 ~~ ~~ a All solutions contain 2 x 1 0-5 mol dmP3 of Fe"PP at pH 10.Wavelength at which maxima appeared in the spectra of the transients, 5 nm. Absolute molar absorption coefficient at these wavelengths, Specific rate constant for the formation of the transient, f 10%. Specific rate constant for decomposition of the transient. f 25 cm3 portions of N,O-saturated solutions containing 4 x lop3 mol dm-, FeI'PP and 0.1 mol dm-, of appropriate alcohol at pH 10.0 were irradiated (with a dose rate 2.5 x lo5 rad hpl) for 1 h.10%. I i0 550 650 i / n m Fig. 2. Spectra of products observed after a pulse producing 3.1 pmol dmP3 of free radicals (N,O-saturated solution at pH 10.0, 2 x lop5 mol dm-, FeIIPP, 0.1 mol dm-3 CH,CH,OH, 12.5 cm optical path): 0, 500 ps after the pulse and A, 20 ms after the pulse.238 REDUCTION OF IRON@) PROTOPORPHYRIN the rate being proportional to [FeIIPP], and the specific rates of reaction thus determined are shown in table. 1. The spectra of the short-lived transients obtained in these reactions, see for example fig. 2, resemble each other but differ considerably from that of FeIPP or FeIIIPP-H. The main difference is in the appearance of a second absorption band with a maximum between 660 and 680 nm. These short-lived intermediates decompose in processes obeying first-order rate laws, the rates being independent of the solute concentrations, pulse intensity and wavelength.The specific rate constants of these reactions are approximately one order of magnitude less than klo. The spectrum of the final product in each of the three systems is identical to that of FeIIIPP. As it has been shown that many aliphatic free radicals react with iron(1r) porphyrins by first forming metal-carbon bonds we suggest that the short-lived intermediates observed are similar, i.e. the first process observed is FerlPP+ 'CR,R,OH -+ Fe1I1PP--CR1R,OH. (12) As the decomposition rate constants of these intermediates are < klo, two plausible mechanisms could be envisaged : (a) heterolytic bond cleavage H** FelllPP-CRIR,OH - FelIIPP+ HCR,R,OH (13) or (6) a redox reaction, as observed for many complexes bound to : CR,R,OH residues H*O Fe111PP-CR,R20H - FelPP+ CRIR,O + H+ (14) followed by reaction (lo), where reaction (14) is the rate-determining step.In order to check these possibilities the yield of dihydrogen in N,O-saturated solutions containing 4 x lop3 mol dmF3 FeIIPP and 0.1 mol dm-3 alcohol at pH 10.0 was measured. The results, shown in table 1, clearly indicate that G(H,) = GH2 + GH only, i.e. no dihydrogen is formed via reaction (14). Thus the results clearly indicate that the decomposition of FeI1IPP-CR1R,OH complexes occurs via reaction (1 3). Note that the rate of this reaction is independent, within experimental accuracy, of the nature of the alcohol.SPECTRA OF THE UNSTABLE INTERMEDIATES Note that the spectra of the three complexes FelllPP-CRIR,OH are very similar to each other. On the other hand the spectrum of FeIIIDP-CH, (where DP = deuteroporphyrin) has only one of the two absorption bands observed for FeIIIPP-CR,R,OH, the band at ca. 570 nm.9a Thus the results indicate that the band at 660-680 nm is due to the a-hydroxy group. The similarity of the spectrum of the intermediate formed in the reduction of Fe"PP by eiq, reaction (8), to that of FeIIIDP-CH, is the reason why we suspect that the transient observed is FeIIIPP-H, formed in reaction (9), and not FeIPP as expected.* We thank Mr D. Carmi for technical assistance and Mr S. Melloul for technical assistance in determining the H, yield in the irradiated solutions.This study was supported in part by the Israel-U.S. Binational Science Foundation (B.S. F.), Jerusalem, Israel. * A final answer to this question can be obtained only by measuring the change in the specific conductivity caused by reaction (8) or (8) + (9). However, we had no access to such an experimental setup.Y. SOREK, H. COHEN AND D . MEYERSTEIN 239 J. E. Falk, Porphyrins and Metalloporphyrins (Elsevier, Amsterdam, 1964). B. Chance, R. W. Estabrook and T. Yonetani, Hemes andHemoproteins (Academic Press, New York, 1966). R. H. Felton and H. Linschitz, J. Am. Chem. Soc., 1966, 88, 143. D. W. Clack and N. S. Hush, J. Am. Chem. Soc., 1968, 90, 4328. D. G. Davis and R. F. Martin, J . Am. Chem. SOC., 1966,88, 1365. T. M. Bednarski and J. Jordan, J.Am. Chem. Soc., 1967, 89, 1552. ' C. Bartocci, F. Scandola, A. Ferri and V. Carassiti, J. Am. Chem. Soc., 1980, 102, 7068. * L. A. Bottomley, L. Olson and K. M. Kadish, in Electrochemical and Spectrochemical Studies of Biological Redox Components, ed. M. Kadish (American Chemical Society, Washington D.C., 1982), p. 219. (a) D. Brault and P. Neta, J. Am. Chem. Soc., 1981, 103, 2705; (b) D. Lexa, J. Mispelter and J. M. Saveant, J. Am. Chem. Soc., 1981, 103, 6806; (c) D. Mansuy and J. P. Battioni, J. Chem. Soc., Chem. Commun., 1982, 638; ( d ) D. Brault and P. Neta, J. Phys. Chem., 1982,86, 3405. l o B. B. Hassinoff and I. Pecht, Biochim. Biophys., 1983, 310. l1 D. Brault, R. Santus, E. J. Land and A. J. Swallow, cited as ref. (31) in ref. (9a). l 2 H. Cohen and D. Meyerstein, Inorg. Chem., 1974, 13, 2434. l3 M. S. Matheson and L. M. Dorfman, Pulse Radiolysis (MIT Press, Cambridge MA, 1969). l4 L. M. Dorfman and G. E. Adams, Nut1 Bur. Stand. Ref Data Ser., 1976, 46. l 5 M. Anbar, Farhataziz and A. B. Ross, Natl Bur. Stand. Ref Data Ser., 1975, 51. l6 M. Anbar, M. Banbeneck and A. B. Ross, Natl Bur. Stand. Ref. Data Ser., 1973,43. l7 J. Butler and A. Henglein, Radiat. Phys. Chem., 1980, 15, 603. l9 P. Neta, Chem. Rev., 1972, 73, 533. 2o (a) G. G. Jayson, J. P. Keene, D. A. Stirling and A. J. Swallow, Trans. Faraday Soc., 1969,65,2453; 21 A. C. Maehly and A. Akeson, Acta Chem. Scand., 1958, 12, 1259. M. Anbar, D. Meyerstein and P. Neta, J. Chem. Soc. B, 1966, 742. (6) J. Halpern, G. Czapski, J. Jortner and G. Stein, Nature (London), 1960, 186, 629. (PAPER 4/772)

ISSN:0300-9599    

DOI:10.1039/F19858100233

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. 1, 1985, 81, 241-243 Comments on Gill’s Approach to the Evaluation of Single Limiting Ionic Conductances in Organic Solvents BY BORIS S. KRUMGALZ* AND ZHENIA FLEISHER Israel National Oceanographic Institute, Tel-Shikmona, Haifa 3 1080, Israel Received 16th May, 1984 Gill’s approach to the evaluation of single ionic conductances in organic solvents is discussed and tested for R,N+ (R = Et, Pr and Bu) and BPh, ions in 11 solvents for which precise transference numbers are known. The results of the test show that the examined method gives comparatively good results (within the average uncertainty declared by Gill and coauthors) only for 3 pure solvents : butan- 1-01, acetonitrile and nitromethane. The problem of splitting the molal function into ionic contributions is not new and has attracted the attention of many researchers for more than 60 years.At the present time, the limiting equivalent conductances can be split experimentally only in some particular cases where experimental transference-number values at infinite dilution have been determined. Since the experimental procedure for transference-number determination in organic solvents is very complicatedl and time consuming, such determinations have been conducted only for a few solutions. Therefore it is necessary to develop an indirect method or methods for splitting the limiting equivalent conductances into ionic contributions in the cases where experimentally determined transference numbers are lacking. The latest review considering all existing indirect methods, beginning from the first approach of Walden et a1.,2 has been published by one of US.^ In this review, a wide variety of assumptions used for developing various indirect methods have been analysed.However, we feel that the last approach developed by Gill4 did not receive proper consideration in our review and requires additional discussion, since this approach has been widely used by Gill and coworkers during recent The equation suggested by Gill and c o ~ o r k e r s ~ - ~ ~ for the determination of 3Lo,i values (for i = Et,N+, Pr4N+, Bu,N+ and BPh;) is given below: where z is the ionic charge, F and N are Faraday’s constant and Avogadro’s number, respectively, qo and E are the solvent viscosity and the relative solvent permittivity, respectively, ri is the ‘true solvated radius’ of the ion under discussion, according to Gill’s definition,3 and r y is an adjustable parameter depending on the solvent structure. Both ri and rY are in A.A t first, Gil14q5 worked out his approach for Et4N+ and only later extended this approach to Pr4N+ and Bu,N+ cations6 and to the BPhy anion.7 Gill4 divided all organic solvents into two groups with characteristic r y values: ( a ) a first group consisting of the highly hydrogen-bonded alcohols and other highly associated solvents having dipolar interactions and (6) a second group consisting of solvents having very little or no association. For the first group of solvents the value ry = 1.13 A 9 24 1 F A R 1Table 1. Comparison of the values calculated using eqn (1) with experimental values N R 25 a - I4 F2 b/Q2-1 cm2 equiv-' "* - 67tNv,[ri - (0.01 0 3 ~ + r,)] no.solvent &25 q;5/10-1 P a s r , / A Et,N+ Pr,N+ Bu,N+ BPh, Et,Nf Pr,N+ Bu,N+ BPh; Et,Nf Pr4N+ Bu,N+ BPh, 1 methanol 32.6c 0.00545d 1.13 59.37 48.01 42.57 38.74 61.12 46.13 39.14 36.47 2 ethanol 24.3e 0.01096 1.13 28.56 23.24 20.67 18.85 28.73 22.55 19.44 - 3 propan- l-ol 20.45s 0.019 67h 1.13 15.68 12.79 11.39 10.40 15.07 12.35 10.89 - 4 butan- l-ol 17.45i 0.026 7f 0.85 10.34 8.60 7.74 7.11 10.20 8.56 7.80 - 1.13 11.42 9.33 8.32 7.60 10.20 8.56 7.80 - 5 acetonitrile 36.G 0.003 44k 0.85 85.77 70.54 63.07 57.72 85.23 70.39 61.70 57.93 6 formamidel 108.7m 0.033 On 0.85 12.24 9.45 8.20 7.35 10.44 7.86 6.54 - 1.13 14.20 10.57 9.03 8.01 10.44 7.86 6.54 - 7 dimethylformamide 37.1m 0.007 93' 0.85 37.36 30.70 27.44 25.11 36.37 29.72 26.68 - 8 dimethylacetamide 37.78p 0.009 19Q 0.85 32.32 26.55 23.73 21.71 32.96 26.2 22.9 - 9 1,1,3,3-tetramethylurea 23.4S 0.014 Olr 0.85 20.12 16.68 14.97 13.74 22.04 17.00 15.47 - 1.13 22.26 18.13 16.13 14.71 22.04 17.00 15.47 - 10 dimethylsulphoxide 46.6s 0.019 63S,t 0.85 15.65 12.77 11.38 10.39 16.39 - 10.93 - 11 nitromethane 35.94c 0.006 2OU 0.85 47.58 39.13 34.99 32.03 47.65 39.14 34.06 33.12 + 2.86 +0.59 - 4.05 - 1.37 - 11.92 - 0.63 - 17.24 - 35.98 - 2.72 + 1.94 +8.71 - 1.01 +4.51 +0.15 - 4.06 - 3.05 - 3.56 - 0.47 - 9.04 -0.21 - 20.23 - 34.50 - 3.30 - 1.34 + 1.88 - 6.63 - + 0.03 -8.77 -6.21 -6.32 - -4.59 - +0.77 - -6.67 - -2.22 0.53 -25.38 - -38.14 - -2.85 - -3.62 - +3.21 - -4.26 ~ -4.12 - -2.73 3.31 a For the calculation of the A::i values from eqn (l), ri has been taken as suggested by Gill,,-' namely rEtaN+ = 4.0, rPrdN+ = 4.6, rBu4N+ = 5.0 and rBPh4 = 5.35 A.The IziPexptl values have been calculated by using only precise transference numbers. The transference numbers used have been discussed by us earlier.3 R. S. Miller and R. M. Fuoss, J . Am. Chem. SOC., 1953, 75, 3076. W. H. Lee and R. J. Wheaton, J . Phys. Chem., 1978, 82, 605. J. Barthel, J. C. Justice and R. Wachter, Z. Phys. Chem. (N.F.), 1973, 84, 100. K. R. Srimvasan and R. L. Kay, J. Solution Chem., 1977, 6, 357. Gill5 noted that his model failed for formamide and N-methylformamide and suggested that this model is applicable only for solvents having dielectric constants between 17 and 77.The smoothed values from S. J. Bass, W. I. Nathan, R. M. Meighan and R. H. Cole, J . Phys. Chem., 1964, 68, 509. J. M. McDowall and C. A. Vincent, J . Chem. SOC., Furaduy Trans. I , 1974, 70, 1862. O L. R. Dawson and W. W. Wharton, J. Electrochem. SOC., 1960, 107, 710. P G. R. Leader and J. R. Gormley, J . Am. Chem. SOC., 1951, 73, 5731. Q G. R. Lester, T. A. Gover and P. G. Sears, J. Phys. Chem., 1956, 60, 1076. B. J. Barker and J. A. Caruso, J . Am. Chem. SOC., 1971, 93, 1341. C. Atlani and J. C. Justice, J . Solution Chem., 1975, 4, 955. The average value calculated from the data: R. L. Kay, S. C. Blum and H. I. Schiff, J. Phys. Chem., 1963, 67, 1223; M. A. Coplan and R. M. Fuoss, J. Phys. Chem., 1964, 68, 1181; J. F. Coetzee and G. P. Cunningham, J .Am. Chem. SOC., 1965, 87, 2529; M. A. Coplan, M. C. Justice and M. Quintin, J . Chim. Phys. Phys. Chim. Biol., 1968, 65, 1152. D. S . Gill, J . Solution Chem., 1979, 8, 691. f Our experimental values. D. F. Evans and P. Gardam, J . Phys. Chem., 1968, 72, 3281. D. F. Evans and P. Gardam, J . Phys. Chem., 1969, 73, 158. S. Minc and L. Werblan, Roczn. Chem., 1966, 40, 1537. P. G. Sears, G. R. Lester and L. R. Dawson, J. Phys. Chem., 1956, 60, 1433. z z n AB. S. KRUMGALZ AND Z. FLEISHER 243 and for the second group the value ry = 0.85 A have been used in all his article^.^-^^ ‘These values of the r y parameter have been obtained4 using the values of ri considered as crystallographic radii or ‘ true solvated radii ’ of Et,N+, Pr,N+ and Bu,N+ ions, and their values have been taken from Coetzee and Cunningham13 as rEtJpr-+ = 4.00, rprlN- = 4.60 and rHuaS+ = 5.00 A.The same method for ri calculations has been used also by Ramanathan et ul.14 However, the ri values obtained by these two groups of researchers differ from each other, although they used the same model. A detailed analysis of all data concerning the dimensions of tetra-alkylammonium cations estimated by various methods has been carried out by one of us previo~s1y.l~ The wide spectrum of the ionic dimensions of R,N’ ions obtained by the various authors shows that the ri value in Gill’s approach [eqn (l)] can be considered as a second adjustable parameter. Therefore, eqn (1) consists of two adjustable parameters, ri and ry, depending on the nature of the ions and solvents, respectively.Gill’s approach to the evaluation of ionic limiting conductances can be tested at present only in 11 organic solvents3 for which precise transference numbers have been experimentally determined. The results of such verification are presented in table 1. The analysis of the results presented in table 1 shows that the differences between the values found experimentally and those calculated by Gill’s approach are from + 8.7 1 7; (for theEt,N+ ion in 1,1,3,3-tetramethylurea) to - 8.77% (for the Bu,N+ ion in methanol) with the exception of formamide solution. As can be seen from table 1, comparatively good results have been obtained only for 3 pure solvents out of 11 tested, namely butan-1 -01, acetonitrile and nitromethane (all with r y = 0.85).P r e v i ~ u s l y ~ - ~ ~ Gill and coworkers declared an average uncertainty of ca. f 2% . If all eleven solvents are discussed, the average uncertainty in lo, calculations is really ca. Ifr 2.9% (without formamide solutions) or f 5.8% (with formamide solutions), but the scattering of the deviations between the experimental and calculated values is extremely large. This fact leads to the conclusion that the use of Gill’s approach to the evaluation of single ionic conductances in organic solvents demands careful testing in each particular case. M. Spiro, in Physical Chemistry of Organic Soltlent Systems, ed. A. K. Covington and T. Dickinson (Plenum Press, New York, 1973), p. 615. P. Walden, H. Ulich and G. Busch, Z. Phys. Chem., 1926, 123, 429. B. S. Krumgalz, J . Chem. Soc., Faraday Trans. I , 1983, 79, 571. D. S. Gill, Electrochim. Acta, 1977, 22, 491. D. S. Gill, Electrochim. Acta, 1979, 24, 701. D. S. Gill, J . Chem. Soc., Faraday Truns. I , 1981, 77, 751. D. S. Gill and M. B. Sekhri, J. Chem. Soc., Faraday Trans. I , 1982, 78, D. S. Gill and J. S. Cheema, Electrochim. Acta, 1982, 27, 1267. D. S. Gill and A. N. Sharma, J . Chem. Soc., Faraday Trans. I , 1982, 78 * D. S. Gill and J. S. Cheema, Electrochim. Acta, 1982, 27, 755. ‘ I D. S. Gill and A. N. Sharma, Ind. J . Chem. A , 1982, 21, 1060. 19. 465. l 2 D. S. Gill, M. S. Chanhan and M. B. Sekhri, J. Chem. Soc., Faraday Trans. I . 1982, 78, 3461. l 4 P. S. Ramanathan, C . V. Krishnan and H. L. Friedman, J . Solution Chem., 1972, 1, 237. l 5 B. S. Krumgalz, J . Chrm. Soc., Faraday Trans. I , 1982, 78, 437. J. F. Coetzee and G. P. Cunningham, J. Am. Chem. Soc., 1965, 87, 2529. (PAPER 4/807) Y-2

ISSN:0300-9599    

DOI:10.1039/F19858100241

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J . Chem. SOC., Furuday Trans. 1, 1985, 81, 245-253 Investigations of the Dynamic Behaviour of Counterions of Anionic Micellar Systems by Fluorescence Quenching Experiments BY MAHMOUD H. ABDEL-KADERP AND AND& M. BRAUN* Institut de Chimie Physique, EPFL-Ecublens, CH- 101 5 Lausanne, Switzerland AND NICOLE PAILLOUS Laboratoire IMRCP, Universite Paul Sabatier, Toulouse 3 1062, France Received 21st May, 1984 The mixing of sodium and nickel lauryl sulphate [SLS and Ni(LS),, respectively] has been investigated using 2-phenylbenzoxazole (1) as a fluorescence probe. Measurements of the fluorescence quenching efficiency above and below the c.m.c. of SLS/or Ni(LS), and analysis of the dependence of the absorption and fluorescence spectra upon solvent polarity lead to the conclusion that the probe is solubilized at the micellar interface and can be used to monitor Ni2+ counterions at the solubilization site.A Stern-Volmer analysis yields non-linear plots reaching a limit of saturation and reflecting the strong adsorption of Ni2+ to the micellar system. Repeating the quenching experiment at constant ionic strength (0.1 mol dm-3) reduces the fluorescence efficiency from 72% to 31% and gives a linear plot. The rate constants of fluorescence quenching have been found to be 4.2 x 1O1O dm3 mol-l s-l when NiSO, is added to SLS micelles and 4.7 x 1Olo dm3 mol-l s-' when Ni(LS), is mixed with SLS. These values are in good agreement with the theoretical diffusion rate constant (3.3 x 1O1O dm3 mol-l s-l) obtained from model calculations using the Debye-Smoluchowski equation.The results obtained fit an intramicellar static quenching process allowing evaluation of the association constant of the Ni2+ counterions with lauryl sulphate micelles. Benzoxazole derivatives are of considerable photochemical and photophysical interest. Whereas the heterocyclic oxazole does no fluoresce itself, it has a powerful fluorescence-enhancing effect when coupled with aromatic systems.' Among these latter compounds, 2-phenylbenzoxazole (1) is of special interest because of its high fluorescence quantum yield and its lasing ability in the U.V. spectral region.2 Its short-lived excited singlet state has led to a number of technical applications such as solar screens3 and stabi1ize1-s.~ The photolysis of 2-phenylbenzoxazole leads to a 2+ 2 dimer.5 Fluorescence quenching is used extensively in studying structural and dynamic aspects of organized assemblies.6 -lo However, results can rarely be generalized as quenching mechanisms differ from one system to another, depending on the fluoro- phore, the quencher and the surfactant used.Fluorescence quenching of 1 has been used in order to study the mixing of sodium and nickel lauryl sulphate [SLS (2) and Ni(LS), (3), respectively]. CH, - [CHJ,, -0 SO; N a* (CH,-KH,I,,- OSO-3),N i *' 2 1 3 t Permanent address: Department of Chemistry, Faculty of Science, Tanta University, Egypt. 245246 DYNAMIC BEHAVIOUR IN MICELLAR SYSTEMS We were thus always dealing with qualitatively similar micellar aggregates but changing their counterions. 1, solubilized near or at the micellar interface, is used as a fluorescent probe in order to monitor the quenching process by the Ni2+ ions at the solubilization site.The mixing of 2 and 3 at concentrations below and above their respective critical micellar concentrations (c.m.c.) should give further insight into the nature of ionic binding of cations of different charges to the same micellar surface. EXPERIMENTAL 2-Phenylbenzoxazole was prepared and purified as reported. l 1 Sodium lauryl sulphate (Fluka, puriss.) has been recrystallized twice from ethanol. Nickel lauryl sulphate was prepared as reported.12 Cyclohexane (Fluka, puriss.) was used as supplied. NiSO, and Na,SO, (Fluka, puriss. p.a.) were used without further purification. Deionized water was distilled twice from a quartz still.Since 1 is sparingly soluble in H,O, a lop5 mol dm-3 solution was prepared from a saturated solution, stirred for two days, by filtering and diluting according to the absorption spectra. Fluorescence quantum yields were measured in SLS micelles using naphthalene as a standard (& = 0.12 in 95% ethanol13 and Lexc = 300 nm at the same optical density). For absorption spectra, a Perkin-Elmer-Hitachi 340 spectrophotometer was used. Fluores- cence measurements were carried out with a Perkin-Elmer MPF-44A spectrofluorimeter with a corrected spectra unit. Fluorescence lifetimes were measured at ca. 25 "C by single-photon counting (Ortec). RESULTS FLUORESCENCE AND ABSORPTION SPECTRA 2-Phenylbenzoxazole, although sparingly soluble in water, will dissolve readily in aqueous SLS micelles.Absorption and emission spectra of mol dmP3 solutions of 1 in water, cyclohexane and SLS micelles are shown in fig. 1 . 1 exhibits an intense absorption band with A, = 300 nm (E,,, = 26.9 x lo3 dm3 mol-1 cm-l in cyclohexane), indicating the n-n* character of the electronic transition to the first excited singlet state. The structure of absorption and fluorescence spectra is solvent dependent. The absorption spectrum in cyclohexane exhibits three major peaks at 292, 300 and 314 nm. The corresponding fluorescence spectrum is a mirror image with maxima at 318, 333 and 350 nm. In H,O the structure of the emission spectrum observed in cyclohexane is replaced by a composite peak. A characteristic feature of the emission spectra in lop2 mol dmP3 aqueous SLS micelles is that a less polar environment of the fluorescent probe is still detectable : the spectra remain slightly structured, indicating that the 2-phenylbenzoxazole molecules are incorporated into the detergent layer of the micelles and at the interface of the aggregates.FLUORESCENCE QUENCHING AT SURFACTANT CONCENTRATIONS BELOW THE C. M.C. For the Stern-Volmer analyses, the lifetime of the fluorescent state of 1 in aqueous solution and in an SLS micellar systems is required. The corresponding results are obtained within the same limits of error to be (7.5k2.5) x 10-lo s. The fluorescence decay is monoexponential in those specific experiments. The quantum yield of the fluorescence of 1 in SLS micelles has been found to agree with that reported for the aqueous solution: #F = 0.78 &- 0.02.'.2a No quenching has been observed in aqueous solutions of 1 and NiSO, at concentrations of up to 5 x An interesting situation arises when the functional surfactant 3 is used as a quencher mol dm-3.M. H. ABDEL-KADER, A. M. BRAUN AND N. PAILLOUS 0 -25 0.20 0.15 0.10 0.05 A/ nm 247 Fig. 1. Absorption and corrected fluorescence spectra of 1 0-5 mol dm-3 2-phenylbenzoxazole at room temperature in different solvents (I = 1 cm and AeX = 300 nm). (-) Cyclohexane, (---) micellar solution (lop2 rnol dm-3 SLS) and (---.) aqueous solution. 0.5 0 5 10 15 20 25 [Ni(LS)2]/10-4 mol dm-3 Fig. 2. Fluorescence quenching of excited 2-phenylbenzoxazole by Ni(LS),. instead of NiSO,. As shown in fig. 2, the ratio of fluorescence intensities ( & / I ) as a function of [Ni2+] does not follow the simple Stern-Volmer analysis : fluorescence quenching is only observed at [Ni2+] > 2 x mol dmP3 and its concentration dependence exhibits a slightly upward curvature up to 6 x lo-, mol dm-3.Addition of 3 results in a further increase of the fluorescence quenching efficiency, until a plateau is reached at [Ni2+] > 1.2 x mol dm-3, the c.m.c. of aqueous systems of 3.14 FLUORESCENCE QUENCHING AT SURFACTANT CONCENTRATIONS ABOVE THE C.M.C. The above results show that the fluorescence quenching of 1 by Ni2+ of 3 is most efficient when quencher concentrations are used permitting micellar aggregation. It248 DYNAMIC BEHAVIOUR IN MICELLAR SYSTEMS 2.5 3. 2.0 e 1.5 1 . o 0.5 0 2 4 6 8 [Ni2’] / rnol dm-3 4 2 1 0 3 6 9 12 [ Ni2+] / 1 Oe2 rnol dm-3 Fig.3. Stern-Volmer analysis of fluorescence quenching of 1 by (a) NiSO, ( lop2 mol dmP3 SLS) and (b) NiSO, ( lop2 mol dm-3 SLS) at constant ionic strength (total salt concentration, 10-l mol drn-”). is of interest to study this quenching behaviour in a case where the fluorophore 1 is already solubilized in an inert micellar system. For this study, two different approaches have been chosen: (a) the fluorophore 1 was solubilized in mol dm-3 aqueous micelles of 2 and NiSO, was used as quencher and (b) 1 was solubilized in mixtures of 2 and 3 of different concentrations, but maintaining a constant concentration of the fluorophore. rnol dm-3) and micellar aggregation number fi (62)15 of 2, the micelle concentration [MI in the absence of any other electrolyte is 2.9 x mol dmp3), multiple occupancy of a micelle by the probe can be avoided.In fact, Poisson statistics16 reveals for the given experimental conditions 71 % empty micelles and 24% singly and 4% doubly occupied micelles. Thus, both self-quenching of the fluorophore and phot~dimerization~ can be neglected. Addition of NiSO, to SLS micelles containing 1 yields a non-linear Stern-Volmer plot [fig. 3 (a)]. In fact, with increasing NiSO, concentration a concomitant increase of the ionic strength of the solution alters the micellar system and thus its concentration. With the given c.m.c. (8.2 x mol dmP3. Using a low concentration of the fluorophoreM. H. ABDEL-KADER, A. M. BRAUN AND N. PAILLOUS 249 Table 1.Surfactant composition, aggregation number and micellar concentration of the system SLS/Ni(LS), [SLSI "i(W21 P I c.m.c. [micelle] no. / 10-3 mol dm-3 / 10-3 mol dmP3 / 10-2 mol dm-3 A / 1 OW3 mol dm-3 / mol dm-3 1 10.0 - 1 .o 62 8.2 0.29 2 8.0 2.0 1.2 68 6.8 0.76 3 6.0 4.0 1.4 73 5.4 1.2 4 4.0 6.0 1.6 79 4.0 1.5 5 2.0 8.0 1.8 84 2.6 1.8 6 - 10.0 2.0 90 1.2 2.0 0 2 4 6 a 10 0.5 ' [ Ni( LS)2 ] / 1 0-3 rnol dm-3 Fig. 4. Stern-Volmer analysis of fluorescence quenching of 1 by Ni(LS), (a) at various mixtures of SLS (lo-, mol dm-3) and Ni(LS), (lo-, mol dmP3), see table 1, and (b) as (a) but maintaining constant ionic strength (0.1 mol dmP3 of Na,SO,). In addition, the strong attraction of the doubly charged quencher cations by the micellar surface alters the statistical distribution of Ni2+.Repeating the quenching experiments at a constant ionic strength (0.1 mol dm-3 by addition of corresponding concentrations of Na,SO,) a linear S tern-Volmer analysis is obtained [fig. 3(b)] from which k, is calculated to be '= 4.2 x 1O1O dm3 mol-l s-l. The second approach to study the fluorescent quenching of 1 in mixtures of 2 and 3 affords the mixing of two solutions of 1 OP5 mol dmP3 of 1 in 1 OP2 mol dmP3 aqueous 2 and mol dm-3 aqueous 3 at various aliquots up to a given volume. The resulting mixed systems all contain the same fluorophore concentration, whereas the micellar concentration [MI differs slightly owing to the different A of 2 and of 3 (table 1). The results show that the quenching process occurs with high efficiency as soon as Ni2+ is present: the plot of & / I = fT[Niz+]) is not linear; maximum efficiency of fluorescence quenching is reached at [Ni2+] 2 4 x On adding Na,SO, (0.1 mol dm-3) to these mixed systems, the fluorescence quenching efficiency drops from 72 to 31 x.However, a linear dependence of I,/I on [Ni2+] is obtained [fig. 4(b)] from which k, is calculated to be 4.7 x 1Olo dm3 mol-1 s-l. mol dmP3 [fig. 4(a)].250 DYNAMIC BEHAVIOUR IN MICELLAR SYSTEMS DISCUSSION The observation that the fluorescence of aqueous solutions of 2-phenylbenzoxazole is not quenched by NiSO, at concentrations up to 5 x mol dm-3 can be easily explained by the short lifetime of the excited fluorophore. The highly efficient quenching obtained in micellar systems, at a much lower concentration of both the quencher and the fluorophore, is also expected.Analogous effects have been reported and rate enhancements in functionalized micellar systems are often attributed to the high local concentration of the reactant adsorbed at or bound to the micellar interface.17 In order to discuss possible mechanisms of fluorescence quenching, the solubilization sites of both the fluorophore and the quencher have to be known. The following experimental facts give conclusive proof that 1 is solubilized at the micellar interface. (1) Absorption and fluorescence spectra of 1 in aqueous micelles of 2 display an intermediate structure of a hydrophilic (H,O) and a hydrophobic (cyclohexane) environment. (2) The non-linear Stern-Volmer plot in fig 4 (a) indicates practically constant and efficient fluorescence quenching at [Ni(LS),] above the c.m.c.Again the fluorescence spectrum of 1 becomes structured allowing aggregates. Our results lead to the conclusion that electronically excited 1 is quenched by the Ni2+ ions associated with micelles and held electrostatically within the Stern as well as the Gouy-Chapman layers. We are thus dealing with an intramicellar quenching process for all the experimental conditions reported. This mechanism can operate in the absence of any type of complex formation in the ground state, which could be detected from absorption or excitation spectra, and has been well illustrated for ‘immobile’ and ‘mobile’ quenchers in nature.ls The non-linear Stern-Volmer plots [fig. 3(a) and 4(a)] reflect very well the adsorption of Ni2+ onto the micellar surface.However, the quenching behaviour in the presence of an electrolyte must still be explained. From fig. 3 and 4 it can be shown that the addition of an inert electrolyte (Na,SO,) decreases the quenching efficiency and consequently the rate of fluorescence quenching. Upon addition of electrolyte, the micellar aggregation number will increase as the repulsion between the charged polar ends of the surfactant is reduced. In fact, it has been found that the aggregation number of sodium cetyl sulphate (SCS) micelles (0.1 rnol dm-3) increases from 67 to 90 when 0.1 mol dm-3 Na,SO, is added.lg Addition of 0.1 mol dm-3 NaCl increases the aggregation number of SLS micelles from 62 to 95,O and decreases the c.m.c.from 8.2 x mol drnP3.,l Thus at comparable quencher concentrations and considering the low probability of even single occupancy by the fluorophore, the fluorescence quenching efficiency has to be lower. However, when the high electrolyte concentration is obtained by the addition of a salt of the quencher (NiSO,), the excess of Ni2+ apparently over-rides these arguments (fig. 3). Note that the aggregation number of 3 is considerably higher than that of 2. This may be due to the stronger adsorption of Ni2+ ions to the micellar surface, which in turn leads to more efficient neutralization of the surface charge. Assuming an aggregation number of 90 for a pure Ni(LS), micellar system14 and a linear dependence of the aggregation number upon dilution of this micellar system by SLS, the values shown in table 1 can be calculated.It follows that at the given point of [Ni2+] = 4 x rnol dmP3 the actual micellar concentration is 1.2 x lop4 mol dm-3 and thus until a ratio of ca. 35 Ni2+ ions per micelle is met all Ni2+ are bound to the interface. Consequently, Ni(LS), systems diluted with SLS to 1.5 xM. H. ABDEL-KADER, A. M. BRAUN AND N . PAILLOUS 25 1 remain unmixed with reference to the counterions until the ratio Na+/Ni2+ exceeds 1.5. Added electrolyte causes a pronounced reduction of the electrical surface poten- tials. It has been reported that the electrical surface potential of SLS micelles of Yo = - 150 mV ( p = 0.008 mol dm-3) is reduced to Yo = -75 mV (11 = 0.1 mol dm-3) upon addition of Na,S04. This causes a decrease of the potential distance function of the Gouy-Chapman layer to approximately one-half of its value.The gradient of the counterion’s concentration in the vicinity of the micelles as described by Boltzmann’s law and calculated from a comprehensive collection of calculated data will follow the variations of the potential distance function.22 Taking into account the low concentration of detergent used, a phase transition from spherical to rod-shaped aggregates on the addition of 0.1 mol dmP3 Na,SO, is not considered. Previous investigations on the dynamics of counterion interactions with charged micellar surfaces23 indicate two possible mechanisms for counterion transfer and the regulation of local concentrations : (a) exchange with the bulk intermicellar phase leading to an equilibriumz4 k+ M f C r M C k- where C is a quencher counterion and M and MC are micellar aggregates containing n and n + 1 quencher counterions, respectively : and (b) transfer of a counterion from the host micelle to another micelle during collision of the two aggregates:25 h-e M,C+M,+M,+M,C.For the exchange mechanism (a), diffusion constants can be evaluated using Debye-Smoluchowski equation :26 k , = MY04) the the where A = 4nNRD/103 is the encounter rate for the neutralized species (spherical micellar aggregate) having a total diffusion coefficient D and an encounter radius R of the micelle plus Ni2+ counterion, N is Avogadro’s number and f(Yo, p ) = (- l / R ) [exp (- I/R) - 13 is a correction term for encounters of charged species which depends on the surface potential of the micelle, Yo, and on the ionic strength p ; 1 is the Onsager distance.Taking for Ni2+ and SLS micelles R = 25 A, I = 45 A and D = 0.8 x lop5 cm2 s-1,24-27 k , is calculated to be 3.3 x 1Olo dm3 mol-l s-l. The theor- etical result is very close to the measured k, of 4.2 x 1O1O dm3 mol-1 s-l when adding NiSO, to SLS micelles in 0.1 mol dm-3 aqueous NaSO, [fig. 3 (b)] and to the measured k , of 4.7 x 1O1O dm3 mol-1 s - l when adding SLS to Ni(LS), micelles in 0.1 mol dm-3 aqueous Na2S0, [fig. 4(6)]. Both results fit very well with k , values previously calculated from experiments (2.5 x 1O1O dm3 mol-1 s-l) with methylene iodide and SLS with pyrene as a f l u o r ~ p h o r e ~ ~ and with calculated values for Cu2+ and SLS micelles (ca.2.9 x 1 O 1 O dm3 mol-l s - ~ ) . ~ , For micelle/micelle collisions, rate constants of I .2 x lox 29 to252 DYNAMIC BEHAVIOUR IN MICELLAR SYSTEMS 4.0 x lo9 dm3 mol-l s-l 24a for probes and quenchers have been 28 We thus deduce that the association/dissociation mechanism (a) is dominant in both experiments when the number of quencher ions remains less than the aggregation number [system Ni(LS),/SLS, table 11 and when Ni2+ ions are added in large excess {[Ni*+]/[M] from 340 to 3400, fig. 3(b), system NiSO,/SLSf. In conclusion, evidence for static quenching in a micellar system can be obtained from experiments where the fluoroscence probe 1 is solubilized at the interface of LS micelles with different concentrations of Na+ and Ni2+ counterions. A quantitative analysis leads to evaluation of the association constant of the Ni2+ ion with the SLS micellar system.The fluorescence quenching experiments for various mixtures of SLS and Ni( LS), were repeated several times in laboratory classes and we thank all the students who showed an interest in this work. The single-photon counting experiments were performed by Dr Robin Humphrey-Baker. We thank Prof. M. Gratzel for permission to use his spectroscopic equipment. M. H. A. thanks the EPFL for a stipend permitting his stay at Lausanne. Financial support by the Swiss National Science Foundation (project no. 4.397-0.80.04) is acknowledged. A. Reiser, L. J. Leyshon, D. Saunders, M. V. Mijovic, A. Bright and J. Bogie, J. Am. Chem. SOC., 1972,94, 2414. * (a)P. J. Roussilheand N.Paillous, J. Chim. Phys., 1983,80,595; (6) C. Rulliere and J. Joussot-Dubien, Rev. Phys. Appl., 1979, 14, 303; (c) C. Rulliere, J. Bellocq, J. Joussot-Dubien and A. T. Balaban, J. Chim. Phys. Chim. Biol., 1978, 75, 961. French Patent 1494097, 1967. (a) G. hick Jr, C. A. Kelly and J. C. Martin, U.S. Patent 4075162, 1978; (b) G. Irick Jr and C. A. Kelly, U.S. Patent 4096115, 1978. J. Roussilhe, E. Fargin, A. Lopez, B. Despax and N. Paillous, J . Org. Chem., 1983, 48, 3736. (a) M. A. J. Rodgers and M. F. Dasilvaewheeler, Chem. Phys. Lett., 1978,53,165; (b) M. A. J. Rodgers and J. C. Becker, J. Phys. Chem., 1980,84,2762; (c) Y. Waka, K. Hamamoto and N. Mataga, Chem. Phys. Lett., 1979, 62, 364; ( d ) A. Yekta, M. Aikawa and N. J. Turro, Chem. Phys. Lett., 1979, 63, 543; (e) D. J.Miller, U. K. A. Klein and M. Hauser, J. Chem. SOC., Faraday Trans. I, 1977,73, 1654. ’ (a) N. J. Turro, M. Gratzel and A. M. Braun, Angew. Chem., 1980, 92, 712; (b) P. P. Infelta, M. Gratzel and J. K. Thomas, J. Phys. Chem., 1974, 78, 190; (c) K. Kalyanasundaram, Chem. SOC. Reti., 1978, 7, 453; ( d ) A. Henglein and R. Scherer, Ber. Bunsenges. Phys. Chem., 1978, 82, 1107. (a) J. C. Russel, D. G. Whitten and A. M. Braun, J. Am. Chem. SOC., 1981, 103, 3219; (b) T. K. Foreman, W. M. Sobold and D. G. Whitten, J. Am. Chem. Soc., 1981, 103, 5333. (a) S. J. Atherton, J. H. Baxendale and B. M. Hoey, J. Chem. SOC., Faraday Trans. I, 1982,78, 2168; (b) M. Almgren, F. Grieser and J. K. Thomas, J. Am. Chem. Soc., 1979, 101, 279; (c) T. Miyashita, T. Murakata and M.Matsuda, J. Phys. Chem., 1983, 87, 4529. lo (a) Y. Croonen, E. Gelade, M. van der Zegel, M. van der Auweraer, H. Vandendriessche, F. C. De Schryver and M. Almgren, J. Phys. Chem., 1983, 87, 1426; (b) M. Almgren, P. Linse, M. van der Auweraer, F. C. De Schryver, E. Gelade and Y. Croonen, J. Phys. Chem., 1984, 88, 289. l1 B. Despax, N. Paillous, A. Lattes and A. Paillous, J. Appl. Polym. Sci., 1982, 27, 225. l2 Y. Moroi, T. Oyama and R. Matuura, J. Colloid Interface Sci., 1971, 60, 103. l3 B. K. Selinger, Aust. J. Chem., 1966, 19, 825. l4 I. Satake, I. Iwamatsu, S. Hosokawa and R. Matuura, Bull. Chem. SOC. Jpn, 1963, 36, 205. l5 J. H. Fendler and E. Fendler, Catalysis in Micellar and Macromolecular Systems (Academic Press, l6 R. C. Dorrance and T. F. Hunter, J. Chem. Soc., Faraday Trans. I, 1974, 70, 1572. l 7 (a) Y. Moroi, A. M. Braun and M. Gratzel, J. Am. Chem. SOC., 1979, 101, 567; (b) M. Maestri. l8 P. P. Infelta, Chem. Phys. Lett., 1979, 61, 88. l9 D. J. Miller, U. K. A. Klein and M. Hauser, Ber. Bunsenges. Phys. Chem., 1980, 84, 1135. 2o (a) J. N. Phillips and K. J. Mysels, J. Phys. Chem., 1955,59, 325; (b) H. F. Huisman, Proc. Kon. Ned. Akad. Wetensch., Ser. B, 1964,67, 388; ( c ) M. F. Emerson and A. Holtzer, J. Phys. Chem., 1965, 69, 3718. New York, 1975). P. P. Infelta and M. Gratzel, J. Chem. Phys., 1978, 69, 1522.M. H. ABDEL-KADER, A. M. BRAUN AND N. PAILLOUS 253 21 E. Matijevic and B. A. Pethica, Trans. Faraday Soc., 1968, 54, 587. 22 M. Gratzel and J. K. Thomas, J. Phys. Chem.. 1974, 78, 2248. 23 (a) D. Stigter, J. Phys. Chem., 1964, 68, 3603; (b) D. Robb, J . Colloid Interface Sci., 1971, 37, 521; (c) J. Oakes, J . Chem. Soc., Faraday Trans. 2, 1973, 69, 1321. 24 (a) F. Grieser and R. Taush-Treml, J . Am. Chem. Soc., 1980, 102, 7258; (b) F. Grieser, Chem. Phys. Lett.. 1981, 83, 59. 25 A. Henglein and Th. Proske, Ber. Bunsenges. Phys. Chem., 1978, 82, 471. 2fi (a) A. J. Frank, M. Gratzel, A. Henglein and E. Janata, Ber. Bunsenges. Phys. Chem., 1978, 547, 80; ( 6 ) Landolt-Bornsrein, Transport Phenomena Z (Springer-Verlag, Berlin, 5th edn, 1969), part 5, p. 63 1. 27 F. Tokiwa and K. Ohki, J. Phys. Chem., 1967, 71, 1343. 2a S. Atik and A. Singer, Chem. Phys. Lett., 1978, 59, 519. 29 M. Griitzel and J. K. Thomas, J. Am. Chem. Soc., 1973, 95, 6885. (PAPER 4/833)

ISSN:0300-9599    

DOI:10.1039/F19858100245

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1985, 81, 255-258 Effects of Solvent on Stacking Interactions A Spectrophotometric Study of Thionine Dimerization in H 2 0 and D 2 0 BY SANDRO L. FORNILI,* GIUSEPPE SGROI AND LIDIA PALUMBO Istituto di Fisica dell’universita and C.N.R.-G.N.S.M., Via Archirafi 36, 1-90123 Palermo, Italy AND VINCENZO Izzol- Istituto Nazionale di Fisica Nucleare, Via Archirafi 36, 1-90123 Palermo, Italy Received 1st June, 1984 Results of a spectrophotometric study at different temperatures of the monomeredimer equilibrium of thionine solutions in H,O and in D,O confirm and extend previous findings, obtained by studying the self-association of Methylene Blue in aqueous solutions, concerning the involvement of solvent rearrangement processes in the stacking of planar dye molecules.Thionine and Methylene Blue are cationic dyes whose planar molecules in aqueous solution self-associate to form face-to-face dimers and higher aggregates with increasing dye c0ncentration.l This makes them and similar dye molecules good model systems for the study of stacking interactions,2 which play a remarkable role in stabilizing helical conformations of nucleic acids.3 Stacking of dye molecules, which causes metachromatic changes in the optical absorption spectra of aqueous solutions of cationic dyes,4 can be investigated by spectrophotometric techniques. A number of interactions can promote ~tacking,~ for instance hydrophobic interactions, which have been shown to provide contributions to the self-association of Methylene Blue6 and which are also appreciably affected by D20-H,O substitution in the ~ o l v e n t .~ Furthermore, the study of the effects of solvent perturbations caused by the presence of some monohydric alcohols on Methylene Blue dimerization has shown a statistically significant linear enthalpy-ntropy compensation with compen- sation temperature T, z 220 K8 These results have been interpreted as providing evidence of a common kind of ‘structural’ alteration of the solvent, as monitored by dye stacking. [ a:;a ] RZN NR, CI- R = CH,, Methylene Blue R = H, thionine Thionine, which differs from Methylene Blue only in that hydrogen-bond donors are present [-NH, groups instead of -N(CH,), groups], enables us to extend the study of the Methylene Blue-water interaction by altering the solute part of the system.In the present work we report the results of a spectrophotometric study of the dimerization of thionine at various temperatures for dilute dye solutions in H 2 0 and in D,O. t Present address: C.N.R.-I.B.S., Via Archirafi 36, 1-90123 Palerrno, Italy. 255256 STACKING OF DYE MOLECULES EXPERIMENTAL Water from a Millipore Super Q sytem and deuterium oxide (99.8% D,O) from Prochem Ltd were used. Thionine chloride, purchased from Merck, was recrystallized three times from 50% ~ a t e r + e t h a n o l , ~ dried overnight at 100 "C and kept in a desiccator. Dye solutions were prepared by diluting concentrated mother solutions and relative concentrations were determined by weighing. Dilutions were performed using plastic syringes (Becton-Dickinson, Ireland), rather than glass pipettes, in order to avoid dye absorption, which could lower the accuracy of the results.Accurate spectrophotometric data were acquired using a microcomputer-based system.'O For each dye concentration absorption spectra were measured successively from the lowest to the highest temperature used. The measurement at the lowest temperature was then repeated in order to check that no irreversible reaction had occurred. Exposure to light during sample preparation was avoided. RESULTS AND DISCUSSION Absorption spectra in the wavelength range 500-700 nm were collected every I nm at five temperatures ( 5 , 10, 15, 20 and 25 "C). The range of thionine concentrations (from 1.2 x to 3.1 x mol was such that only the monomeredimer equilibrium was relevant, according to the criteria discussed in ref.(7). For each temperature, five dilute dye solutions in H,O and five solutions in D,O, at different dye concentrations, were used to evaluate the dimerization equilibrium constants and the monomer and dimer absorption spectra. This was required by the data processing algorithm we a d ~ p t e d , ~ which implements an iterative least-squares best-fitting procedure based on K = (1 - x)/(2cx2) and E,, = EP'X + e$ ( I - x) where K, c and x are the dimerization constant, the dye concentration and the monomer molar fraction, respectively, and cv, EP and c$ are the apparent, monomer and dimer molar extinction coefficients, respectively. In fig. 1 we plot the logarithm of the thionine dimerization constant against the inverse of the absolute temperature for both H,O and D,O dye solutions.Analogous data for Methylene Blue, taken from ref. (7), are shown for comparison. Standard enthalpy and entropy changes for thionine dimerization were evaluated by linear regression according to the relation- ship AGe = - RTln ( K ) . They are reported in table 1 , together with analogous values for Methylene Blue taken from ref. (7). We can now make the following observations. (1) In the temperature range considered, an approximately linear relationship is found between the logarithm of the dimerization constant and the inverse of the absolute temperature. Furthermore, although the standard enthalpy and entropy changes for thionine dimerization (table 1) show a small difference between H,O and D,O, the corresponding best-fit straight lines practically coincide over the temperature range reported in fig.1. (2) Table 1 shows that dimerization processes in H,O for thionine and for Methylene Blue differ by ca. 120% for the entropy change and by ca. 10% for the enthalpy change. Therefore, the substitution in Methylene Blue molecules of -N(CH,), groups by -NH, groups (conversion to thionine) affects mainly the entropy term of the dimerization free energy. These findings parallel the results reported in previous papers of this series concerning the effects on Methylene Blue dimerization of H-D isotopic substitution in the solvent7 and of solvent perturbation by some monohydric alcohols.8 ( 3 ) In the hypothesis that for thionine, as for Methylene Blue,6 dimerizationS.L. FORNILI, G. SGROI, L. PALUMBO AND V. IZZO 10 : I r 9 --. - a E f :*- - 257 - I 1 I I I 5@+ Table 1. Standard enthalpy and entropy changes for dimerization of Methylene Blue thionine H,O - 7.4 k 0.2 - 12.0k0.9 D2O - 7.6 f 0.2 - 12.5k0.9 Methylene H,O - 6.7 k 0.2 -5.4k0.9 Bluec D2O -6.2 f0.2 -3.5k0.5 hionine and a In kcal mol-l (1 cal = 4.184 J); in e.u. (1 e.u.= 1 cal K-'mol-l); data from ref. (7). causes a decrease in the number of solute molecules without releasing solvent molecules, AS* contains a term related to the mixing entropy. This term (ca. 8 e.u. for the dimerization case) has to be substracted from AS* to give the 'unitary' entropy change,l1V l2 ASP, which results from negative contributions (rotational and translational entropy losses and electrostatic entropy change) and from a positive contribution from the dye-solvent interaction.6 For Methylene Blue dimerization, the positive term is large enough to overcome the negative contributions, the net result being a positive value of ASF.63 For thionine AS? is negative (-4 e.u.), which indicates that for thionine a smaller positive entropy change is contributed by dye-water interactions as compared with the Methylene Blue case, since it is reasonable to suppose that in both cases the negative entropy contributions are similar.For the substitution of each methyl group with hydrogen in the Methylene Blue molecules, a loss of ca. 1.65 e.u. occurs. This result is in full agreement with the above-mentioned258 STACKING OF DYE MOLECULES interpretation of previous experimental result^.^.8 * 1 3 s l4 Indeed, the presence of hydrogen-bond donors in the thionine molecules causes a less relevant rearrangement of the solvent molecules upon self-association of the dye molecules, as compared with the Methylene Blue case. We thank Ms A. La Franca and Mr S. Francofonte for technical help. Indirect support from C.R.R.N.S.M. and M .P.I. is acknowledged. D. G. Duff and C. H. Giles, in Water: A Comprehensive Treatise, ed. F. Franks (Plenum Press, New York, 1975), vol. IV, chap. 3. G. G. Hammes, Adv. Protein Chem., 1968, 23, 1. H. Edelhoch and J. C. Osborne Jr, Adv. Protein Chem., 1976, 30, 183. S. F. Mason, J. SOC. Dyers Colour., 1968, 84, 604. E. Coates, J. SOC. Dyers Colour., 1969, 85, 355. P. Mukerjee and A. K. Gosh, J. Am. Chem. SOC., 1970,92, 6419. ' S. L. Fornili, G. Sgroi and V. Izzo, J. Chem. SOC., Faraday Trans. I , 1981, 77, 3049. S. L. Fornili, G. Sgroi and V. Izzo, J. Chem. SOC., Faraday Trans. I , 1983,79, 1085. G. R. Haugen and E. R. Hardwick, J. Phys. Chem., 1963,67, 725. lo S. L. Fornili, J. Phys. E, 1980, 13, 34. l 1 R. W. Gurney, Ionic Processes in Solution (McGraw-Hill, New York, 1953; Dover, New York, 1962). l 2 W. Kaumann, Adv. Protein Chem., 1959, 14, 1. l 3 R. E. Ballard and C. H. Park, J. Chem. SOC. A , 1970, 1340. l4 L. Cordone, A. Cupane aild S. L. Fornili, Biopolymers, 1983, 22, 1677. (PAPER 4/901)

ISSN:0300-9599    

DOI:10.1039/F19858100255

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J . Chern. Sac., Faraday Trans. 1, 1985, 81, 259-263 Rate Constants for the Reactions of Hydroxyl Radicals with Propane and Ethane BY DONALD L. BAULCH,* IAN M. CAMPBELL AND SANDRA M. SAUNDERS Department of Physical Chemistry, The University, Leeds LS2 9JT Received 5th June, 1984 The kinetics of the reaction OH + C,H, --+ H,O + C3H7 have been studied in a discharge-flow system under first-order conditions. The OH radicals were generated by the reaction of H atoms with NO, and the concentration of OH, monitored by resonance fluorescence, was followed as a function of reaction time in a large excess of the alkane. A value of k , = (7.2+ 1.1) x 10, dm3 mol-1 s-l at 295 K was obtained. As a check on the technique the rate constant for the reaction OH + C,H, -+ H,O + C,H, (2) was determined in a similar fashion.The value obtained, k , = (1.61 kO.24) x 10, dm3 mol-l s-l, is in excellent agreement with other literature values. The reaction of OH radicals with propane OH + C,H, -+ H,O + C,H, (1) is of importance in combustion processes and the chemistry of the atmosphere. A recent evaluation1 of the rate constant for reaction (1) for use in atmospheric modelling demonstrates a surprising lack of concordance in the available data at 300 K, which scatter over a range of nearly a factor of 3. Equally surprising is the fact that the reaction has not been studied under first-order conditions in a discharge-flow system, a technique well suited to measuring this rate constant. The only previous discharge-flow study2 used comparable concentrations of propane and hydroxyl radicals and required a mass-spectrometric measurement of the reaction stoichiometry to derive a value of k,.Such measurements are difficult to perform accurately and introduce a corresponding uncertainty into the rate-constant value so obtained. In the present work OH radicals generated in a discharge-flow system by means H+NO,-+OH+NO of the rapid reaction were reacted with an excess of propane sufficiently large to ensure that there were no significant routes of OH removal other than reaction (1). The rate constant, k,, was derived from the change down the flow tube of [OH], monitored by resonance fluorescence. To test the satisfactory operation of the apparatus and technique the well established' rate constant for the reaction OH -t- C,H, -+ H,O + C,M, (2) was also measured.259260 REACTIONS OF OH WITH C,H8 AND C,H, EXPERIMENTAL The apparatus used was a discharge-flow system of conventional design, having a 25 mm i.d. Pyrex flow tube 1 m in length with inlet ports and a sliding double injector for introduction of reagents. At a fixed point at the downstream end of the tube a brass fluorescence cell for monitoring the [OH] was mounted. H atoms, produced by a microwave discharge in a stream of He containing a small quantity ( 1 % ) of molecular H,, were introduced into the flow tube through a sidearm and titrated with NO, introduced through the outer inlet of the dual injector. In this way known concentrations of OH were produced by the fast reaction H+NO, -+ OH +NO, (3) with k , = 8 x 10'" dm3 mo1-l s-' at 298 K.Other reactants (C,H, and C,H,) were admitted to the flow tube through the central inlet of the injector. The distance between the point of admission of NO, and the alkane was sufficient to ensure effective completion of reaction (3) under the conditions used. The [OH] was monitored by means of a conventional OH resonance fluorescence system. The OH emission band centred on 309 nm ( A 2C t X "n) was generated by passing a microwave discharge through a stream of argon saturated with water vapour. The fluorescence emission was detected after passage through a collimator and narrow-band filter (Oriel 5703-10 nm band width) by a photomultiplier (EM1 9789Q) placed at angles to both the OH lamp and flow-tube axes. The photomultiplier signal was fed to a photon counter (Brookdeal-Ortec PCS 5C1).Light traps to minimize scattered radiation were placed opposite the resonance lamp and photomultiplier. The initial OH concentration was obtained from the value of [NO,] at the end-point of the titration of H with NO,. The variation of the fluorescence intensity was studied as a function of OH concentration and, as has been observed by others, it was found to be linear at [OH] < cu. 1 OPH mol dmP3. All measurements were carried out with OH concentrations well below this value. To reduce wall removal of the OH radicals the flow tube and injector were coated with Halocarbon wax (series 15-00> and the interior of the fluorescence cell was coated with Teflon. The first-order rate constant for wall removal of the OH was found to be cu.23 s-l and constant throughout the series of measurements. Preliminary experiments demonstrated the necessity of removing trace impurities from the helium carrier gas (B.O.C., CP grade), which was therefore passed through a silica tube at 900 K packed with copper turnings, followed by a liquid-nitrogen trap. Argon (B.O.C., A grade), H, and C,H, (Matheson, research grade) and C,H, (Matheson, CP grade) were used without further purification. The NO, (Matheson, 99.57,) was further purified by mixing with 0,. leaving to stand overnight, freezing and pumping away the 0,. subjecting the NO, to a series of freeze-thaw cycles and finally taking the middle fraction. Concentrations of reagents were calculated from flow rates measured by means of calibrated capillary flow meters and the mean pressure in the flow tube.Operating conditions were: total pressure cu. 0.3 kPa; temperature 295 2 K , linear flow velocities cu. 15 m s-* ; [OH],, z mol dmP3. RESULTS Since the resonance-fluorescence detection system is newly constructed its perform- ance, and the other procedures, were checked by measuring the rate constant for reaction (2). Its value is well established (table 1) and a recent evaluation1 has recommended k , = ( I .63 k0.24) x IOH dm3 mob ssl at 298 K. In a very large excess of C,H, the decay of [OH] down the flow tube will be given In [OH]/[OH], = - (k' + k , ) t where k , is the first-order rate constant for reaction of OH with the walls, k' = k,[C,H,] and t, the reaction time. is related to the reaction distance downD.L. BAULCH, I. M. CAMPBELL AND S. M. SAUNDERS 26 1 I I 1 I 0 0.01 0.02 0.03 0.04 t l s [C,H,] = 7.6 x lo-@ rnol dmP3 and 0, [C,H,] = 2.8 x lo-’ rnol dmP3. Fig. 1. Plot of in[OH] as a function of reaction time for [OH], = mol dm-3: 0, the flow tube, d, and the linear flow rate, f,>, by t = d/fIJ. Good first-order plots were obtained for ln[OH]/[OH], against d, as illustrated in fig. 1 for two different values of [C,H,]. Values of (k’ + k,) obtained from the least-mean-squares value of the slopes of such plots were plotted against [C,H,] (fig. 2). Within the limits of scatter the plot is linear with an intercept k , = 2243 s-l and a slope k, = (1.61 kO.10) x lo8 dm3 mol-l s-l, where the error limits are one standard deviation.This value of k, is in excellent agreement with previous measurements and the value of k , agrees closely with the value of 2 3 k 3 s-l measured in the absence of C,H,. mol dm-3 and although pseudo-first-order kinetic behaviour still appeared to apply, the value of k was based only on measurements at [C,H,] > 1.5 x lop7 rnol dm-3 (i.e. [C,H,]/[OH], > 100) under which conditions only reaction (2) and wall removal of OH should be significant. Reaction (1) was studied in a similar fashion. Values of (k’ + k,) as a function of [C3H,] are shown in fig. 2. A least-squares fit yields k , = (7.2 f 1.1) x lo8 dm3 mol-1 s-l. A detailed analysis of sources of error arising in flow-tube studies33 suggest that an overall error of & 150/;, as quoted here would be more realistic than the 804 standard deviation given by the least-squares fit.Although measurements were made for [C,H,] < 1.5 x262 100 8 0 - I co . h e3 60 5 + 40 20 REACTIONS OF OH WITH C,H8 AND C,H, I I I I 0 1.0 2.0 3.0 L.0 [alkane]/lO-' mol dm-3 Fig. 2. Plot of (k'+k,) as a function of alkane concentration. ., C,H6 and D, C3H,. The lines are least-mean-squares fit to the points. Table 1. Experimentally determined values of k , and k, at temperatures in the region of 300 K ~ ~ ~ _ _ _ ~ ~ _ _ _ alkane k/lOs dm3 mol-' s-' techniquea TIK ref. 'ZH6 I .70 1.59 k0. 10 1.74 k 0.36 1.56k0.24 1.39 k 0.24 1.55 k0.13 1.61 kO.10 7.24 & 0.42 5.00k0.12 6.32 & 0.24 9.09 0.13 7.35 & 0.30 12.16 k 0.60 13.24f3.6 7.2 0.60 f. p .-p . p. f. p .-r .a. d. f.--1.m .r. d.f.-r.f. d.f.-ref. f.p.-r.f. d. f.-r . f. d.f.-e.s.r. f.p.-r.f. smog chamber smog chamber f. p.-r . a. photolysis-product analysis p.f. - r.f. f.p.-k.s. 297 295 k 2 296 298 RTb 298 295 &- 2 297 298 298 300 299 295 k 2 298 295 f 2 5 6 7 8 9 10 this work 5 2 10 11 12 6 13 this work a f.p., flash photolysis; p.p., plate photometry; r.a., resonance absorption; d.f., discharge Room flow; l.m.r., laser magnetic resonance; r.f., resonance fluorescence; k.s., kinetic studies. temperature. DISCUSSION The good agreement between our value for k , and other literature data suggest that the apparatus and procedures are satisfactory. The available values of k , at temperatures close to 300 K are given in table 1. The results from the present study agree closely with the most recent, and probably theD .L. BAULCH, I. M. CAMPBELL AND S. M. SAUNDERS 263 most reliable, flash-photolysis work of Tully et a1.lo and also with the earlier flash- photolysis study of Greiner5 and the relative rate measurements of Darnall et a1.l1 The earlier discharge-flow work gives a low value, probably because of the difficulty of determining the stoichiometry with any accuracy. The other results are all very high, although only one of them, that of Overend et aZ.,6 is clearly well outside the quoted error limits; there is no obvious reason for such a discrepancy since the same group obtained an acceptable value for k , using the same technique. On the basis of the present work and that of G r e i n e ~ , ~ Tully et aLlo and Darnall et al." we suggest a value of k, = (7.0 & 0.6) x lo8 dm3 mol-1 s-l at 298 K.We thank the S.E.R.C. for the award of a studentship to S.M.S. during the tenure of which this work was carried out. * CODATA Task Group on Chemical Kinetics, D. L. Baulch, R. A. Cox, R. F. Hampson, J. A. Kerr, J. Troe and R. T. Watson, Evaluated Kinetic and Photochemical Data for Atmospheric Modelling, Supplement II, J. Phys. Chem. ReJ Data, in press. J. N. Bradley, W. Hack, K. Hoyermann and H. Gg. Wagner, J . Chem. SOC., Faraday Trans. 1 , 1973, 69, 1889. C. J. Howard, J. Phys. Chem., 1979, 83, 3. F. Kaufman, Prog. React. Kinet., 1961, 1, 3. N. R. Greiner, J. Chem. Phys., 1970, 53, 1070. R. P. Overend, G. Paraskevopoulos and R. J. Cvetanovic, Can. J. Chem., 1975, 53, 3374. C. J. Howard and K. M. Evenson, J. Chem. Phys., 1976,64,4303. M. T. Leu, J. Chem. Phys., 1979, 70, 1662. J. H. Lee and I. N. Tang, J. Chem. Phys., 1982,77,4459. lo F. P. Tully, A. R. Ravishankara and K. Carr, Int. J . Chem. Kinet., 1983, 15, I 1 1 1 . I1 K. R. Darnall, R. Atkinson and J. N. Pitts, J . Phys. Chem., 1978, 82, 1581. l2 R. Atkinson, S. M. Aschmann, W. P. L. Carter, A. M. Winer and J. N. Pitts, Int. J . Chem. Kinet., l 3 R. A. Gorse and D. H. Volman, J. Photochem., 1974, 3, 1 1 5. 1982, 14, 781. (PAPER 4/924)

ISSN:0300-9599    

DOI:10.1039/F19858100259

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. 1, 1985, 81, 265-266 Molten-salt Hydrate Media Cobalt(I1) Ions in ZnCl, + H,O and CaC1, + H,O Systems BY JOHN A. DUFFY* AND G. LESLIE WOOD Department of Chemistry, The University, Aberdeen AB9 2UE Receiced 5th June, 1984 The solvent properties of the ZnC1, + H,O system towards Co2+ differ from those of other metal chloride aqueous systems. In CaC1, .xH,O, for example, tetrahedral chlorocobaltate(I1) species are increasingly generated with decreasing water content for x = 12, but not in ZnC1,. xH,O, where octahedral coordination of Co2+ is preserved even for x = 3. The importance of very concentrated aqueous solutions of electrolytes has been recognised' and the properties of molten-salt hydrates have been investigated occasionally; for example, the solvent behaviour of the melts MgCl, - 6H,O and CaCl, - 6H,O towards nickel(II).2y We are conducting experiments in aqueous solutions with high concentrations of simple salts, and although our principal concern is with the decomposition of inorganic complexes [chiefly those of cobalt(rx1) and chromium- (III)], we wish to report the marked difference in solvent behaviour between the two molten-salt hydrate systems, CaC1, .xH,O and ZnC1, .xH,O.The difference is demonstrated in a striking manner by the Co2+ ion. In aqueous solution Co2+ exists as the pink octahedral aquo complex, and addition of chloride ions converts this into the blue tetrahedral chlorocobaltate(I1) complex. The conversion is preceded, at lower chloride concentrations, by enhanced absorbance of the octahedral species owing to some replacement of H,O ligands by C1-.These effects are observed in many chloride + water systems and fig. 1 (a) shows data for aqueous solutions of CaCl, which, at very high chloride concentration, exist as the molten-salt hydrate: CaC1;6H2O (m.p. < 30 "C). The contrasting behaviour of Co2+ in the ZnC1, + H,O system is shown in fig. 1 (6). It can be seen that octahedral coordination of the Co2+ is maintained, even in the hydrates ZnC1, .6H,O and ZnC1, 3H,O, and enhanced optical absorption occurs only at these very high chloride concentrations and not at lower. Virtually no tetrahedral chlorocobaltate(r1) species is produced, and this suggests that, unlike the CaCl, system, there are insufficient chlorides available for complexation with the CO".Zinc chloride has a silica-like network structure in the molten state which requires all the chloride ions to be bridging, and presumably in very concentrated aqueous solutions this structure, or a related one, is preserved. (Even in dilute solution, ,Zn<E:>Zn., unitsexist.4) In comparing the two molten- salt hydrate systems, the CaCl,+H,O system behaves in a conventional manner as a concentrated solution of metal ions and anions, whereas the ZnC1, + H,O system might be described as a 'diluted molten salt'. This behaviour fits into the pattern previously observed5* for ZnC1, when the dilution is with molten alkali-metal chlorides. \ / We thank B.P. Petroleum Development Ltd for financial support for this work. 265266 MOLTEN-SALT HYDRATE MEDIA 1 water/salt ratio (x) 500 400 0" 4 300 e 100 0 200 g Fig. 1. Octahedral (0) and tetrahedral (A) species of Co2+ in (a) the CaCl, + H,O system and (b) the ZnCl,+H,O system at ambient temperature. The Co2+ species are identified from their 4T,g(F) -+ 4T,g(P) and 4A,(F) + 4T,(P) absorption bands. The left-hand side absorbance scale refers to octahedral species and the right-hand side scale to tetrahedral species. Ionic Liquids, ed. D. Inman and D. G. Lovering (Plenum Press, New York, 1981). C. A. Angell, J. Phys. Chem., 1965, 69, 2137. D. G. Lovering, Collect. Czech. Chem. Commun., 1972, 37, 3697. K. Sawada, M. Okazaki, H. Hirota and M. Tanaka, Znorg. Chem., 1976, 15, 1976. C. A. Angell and D. M. Gruen, J. Phys. Chem., 1966,70, 1601. W. E. Smith, J. Brynestad and G. P. Smith, J. Am. Chem. SOC., 1967, 89, 5983. (PAPER 4/925)

ISSN:0300-9599    

DOI:10.1039/F19858100265

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J Chem. Soc., Faraduy Trans. I , 1985, 81, 267-271 Conversion of Carbon Monoxide into Methanol at Room Temperature and Atmospheric Pressure BY KOTAROGURA* AND SHINJI YAMASAKI Department of Applied Chemistry, Yamaguchi University, Ube 755, Japan Received 12th June. 1984 It has been found that CO can be converted selectively into methanol using Everitt's salt (ES) in the presence of pentacyanoferrate(I1) and methanol. The formation of methanol was confirmed from isotope distributions analysed by gas chromatography coupled with mass spectrometry. The reduction of CO into methanol is brought about by the oxidation of ES to the Prussian blue (PB). This redox reaction, however, needs to be activated by homogeneous catalysts consisting of hexacyanoferrate(r1) and methanol. The conversion continues to proceed with the consumption of proton in solution under controlled potential conditions such that PB is reduced to ES.Abundant but relatively inert molecules such as nitrogen, carbon monoxide, carbon dioxide and methane are enticing as chemical feedstocks. Carbon monoxide conversion has been investigated in methanol synthesis2 and in Fischer-Tropsh reaction^,^ which essentially proceed over catalysts at the pressure of the synthesis gas near 300 atmt and at temperatures near 200 "C. These severe conditions are unfavourable for the extensive use of CO as a chemical feedstock. Some investigator^^-^ have attempted the reduction of CO under mild conditions. However, their hydrogenation of CO is only or CO conversion gives only squarate anions in non-aqueous solvent^^-^ and dimeric species in liquid a m r n ~ n i a .~ We have carried out selective conversion of CO into methanol with Everitt's salt using pentacyanoferrate(I1) and methanol as homogeneous catalysts at room temperature and atmospheric pressure. Herein we report the results of these studies and discuss the mechanism of conversion of CO into methanol under the present conditions. EXPERIMENTAL Everitt's salt [ES, K,FellFell(CN),] was prepared by the electrochemical reduction of Prussian blue [PB, KFelllFell(CN),], which was electroplated on a 4.1 cm2 platinum plate from a fresh solution of 0.01 mol dm-, FeCl, and 0.01 mol dmP3 K,Fe(CN),.lO-lz The amount of E:S on the platinum plate was evaluated coulometrically in 0.1 mol dm-" KC1. The metal complexes used were aquapentacyanoferrate(r1) {Na,[Fe(CN),(H,O)]j, pentacyanonitrosyl- fi:rrate(Ilr) {Na,[Fe(CN),(NO)]) and amminepentacyanoferrate(H) (Na,,[Fe(CN),(NH,)]>. Aquapentacyanoferrate(1r) was synthesized by the method of Hieber et af.,13 but pentacyanonitrosylferrate(Ii1) and amminepentacyanoferrate(i1) (Katayama, reagent grade) were used as received.Reagent-grade methyl formate or potassium iodide was occasionally used as an additive. Carbon monoxide (Seitetsu Kagaku) was 99.9% pure and had the following impurities: 500 pprn nitrogen, 100 ppm carbon dioxide, 100 ppm oxygen, 100 ppm hydrogen and 15 ppm water. The catalyst solutions were prepared by dissolving given amounts of iron complexes and/or methanol into 0.1 mol dm-" KCI solution (the pH was adjusted with HCI) and kept in a 1 dm3 + 1 atm = 101 325 Pa.267268 CONVERSION OF co INTO METHANOL reservoir into which CO was fed at a flow rate of 100 cm3 min-' for 1 h. In the analysis of isotope distributions in methanol both added and produced, the solution was prepared with D,O (Merck, 99.75%) instead of H,O and the pH was adjusted with DCI (Merck, 20% in D,O). 35 cm3 of the prepared solution were transferred to the test cell which was connected to the counter-electrode (platinum) cell through a fine frit and to the reference-electrode (saturated calomel) cell through a solution-lubricated glass stopcock. A platinum plate coated with PB film was put into the test cell containing 0.1 mol dmP3 KCl and the electrode was polarized to -0.4 V. After the PB film had been reduced to the colourless ES the solution was replaced by the catalyst solution under a nitrogen atmosphere.The experiments were performed under both open-circuit and cathodically controlled potential conditions. Isotope distributions in added and produced methanol were analysed by means of a mass spectrometer (JMS-D100, Nippon Densi Co.). A JGC-1 100-type gas chromatograph with a thermal-conductivity detector and a Porapak Q column was used for the determination of methanol. The sampling procedure was performed by the following way. 2.5 cm3 of the sample solution were transferred to a cell which was connected to a vacuum through a stopcock. A side-port of the cell was fitted with a rubber septum in order to withdraw samples uia a syringe, and the solution was evaporated under 1 Torrt pressure at 70 "C.After nitrogen gas had been introduced into the cell, gas samples (2cm3) were taken with a Pressure Lock air-tight gas syringe. The calibration curve for this sampling procedure was linear for methanol concentrations ranging from 0.01 to 30 mmol dmP3. The amount of methanol produced was calculated from the difference in methanol contents of the initial and final solutions. Formic acid and formaldehyde were determined by a colorimetric analysis using chromotropic acid. l4 The absorption spectrum of the solution was obtained by means of a double-beam spectrometer (Hitachi model 100-50). RESULTS AND DISCUSSION Mass spectra of methanol in a 0.1 mol dm-, KCl+D,O solution containing methanol and pentacyanoferrate(I1) were obtained before and after electrolysis at - 0.9 V.Two peaks were seen at m/e 3 1 and 32 corresponding to CH,O and CH,OH, and these are due to CH,OH added initially. After electrolysis, however, three additional peaks appeared at m/e 33, 34 and 36 which are ascribed to CH,OD, CD,O and CD,OD, respectively. These results show the formation of methanol in the present system. In the progress of the conversion of CO under the open-circuit conditions, the colourless platinum surface coated with ES became intensely blue, showing that ES was oxidized to PB by the reactiong ES 4 PB+K++e-. (1) The reduction of CO into methanol should thus be coupled with this oxidation. The rate ( u ) of methanol formation in this conversion system was linearly related to pH, and log u/pH was close to the value of -0.25/pH. Hence the net process for the reduction can be represented as (2) This reaction can be only activated by a homogeneous catalyst system consisting of a metal complex and methanol.The results of the quantitative analysis of the reaction products under open-circuit condition are given in table 1. In the presence of Fe(CN):- or CH,OH alone (run 1 or 2) there is no formation of methanol. Neither is it obtained in the absence of ES (run 5). Therefore the coexistence of Fe(CN);-, CH,OH and ES (runs 3 and 4) is essential for the conversion of CO into methanol. Methanol formation is very selective, since the amounts of formaldehyde and formic acid are negligibly small. The CO + 4H+ + 4ES + CH,OH + 4PB + 4K+. t I Torr = 101 325/760 Pa.K.OGURA AND S. YAMASAKI 269 Table 1. Results of the quantitative analysis of the reaction products under open-circuit conditions in 0.1 mol dmP3 KCla [CH,OH] [HCHO] [HCOOH] (ES)" run catalystb / 1 Op7 mol cm-, /pmol dm-, 1 A 2.92 0.0 0.5 1 0.38 2 B 2.94 0.0 0.52 0.20 3 C 2.67 600 0.42 0.36 4 D 2.93 1190 0.63 1.02 5 D 0.00 0.0 0.02 0.03 6 E 2.73 I700 0.08 25.03 7 F 2.60 1710 0.46 0.49 Reaction time 3 h, temperature 23 OC, pressure I atm and initial pH 3.5. (A) 10 mmol- dmP3 CH,OH, (B) 5 mmol dm-, Na,[Fe(CN),(H,O)], (C) 5 mmol dm-, Na,[Fe(CN),(H,O)] + 10 mmol dm-, CH,OH, (D) 5 mmol dm-, Na,[Fe(CN),(H,O)] + 20 mmol dm-, CH,OH, (E) 5 mmol dm-, Na,[Fe(CN),(H,O)] + 20 mmol dm-, CH,OH + 10 mmol dm-, HCOOCH, and (F) 5 mmol dm-, Na,[Fe(CN),(H,O)] + 20 mmol dm-, CH,OH + 10 mmol dm-, KI.The amount of Everitt's salt coated on Pt. Table 2. Results of the quantitative analysis of the reaction products under cathodically polarized conditions in 0.1 mol dm-, KCla [CH,OH] [HCHO] [HCOOH] (W" run catalystb /lo-' mol cmP2 /pmol dm-, 1 2 3 4 5 6 7 8 9 10 1 I d A B C D D E F G H 1 J 3.1 1 2.72 2.36 2.22 0.00 2.95 2.64 2.18 3.75 2.85 2.93 0.0 0.0 2180 3490 4560 4360 0.0 0.0, 0.03 2460 290 1.46 0.14 0.46 0.79 1.01 0.08 0.04 0.08 0.02 0.33 0.12 0.0 1 0.19 0.77 0.35 0.02 8.98 7.09 0.59 5.49 0.00 0.0 1 Cathodizing potential -0.9 V (27s SCE), reaction time 3 h, temperature 23 "C, pressure 1 atm and initial pH 3.5. (A) 10 mmol dm-" CH,OH, (B) 5 mmol dm-, Na,[Fe(CN),(H,O)], (C) 5 mmol dm-, Na,[Fe(CN),(H,O)]+ 10 mmol dm-, CH,OH, (D) 5 mmol dm-, Na,[Fe- (CN),(H,O)] + 20 mmol dm-, CH,OH, (E) 5 mmol dm-, Na,[Fe (CN),(H,O)] + 20 mmol- dm-3 CH,OH + 10 mmol dmP3 HCOOCH,, (F) 5 mmol dmP3 Na,[Fe(CN),(H,O)] + 20 mmol dm-, CH,OH + 10 mmol dm-3 KI, (G) 5 mmol dmP3 K,[Fe(CN),] + 20 mmol- dm-, CH,OH, (H) 5 mmol dm-, Na,[Fe(CN),(NO)] + 20 mmol dm-, CH,OH, (1) 5 m m ~ l d m - ~ Na,[Fe(CN),(NH,)] + 20 mmol dmP3 CH,OH and (J) 5 mmol dm-, Na,[Fe(CN),(H,O)] + 20 mmol dm-, HCOOCH,.The catalyst solution was not bubbled with CO. The amount of Everitt's salt coated on Pt.270 CONVERSION OF co INTO METHANOL addition of methyl formate (run 6) or potassium iodide (run 7) enhances the formation of methanol. The function of these additives appears to be related to the stabilization of a reaction intermediate as discussed latter.Similar results under cathodically controlled conditions are shown in table 2. As in table I , the coexistence of cyanoferrate(II), methanol and ES is necessary for CO conversion, and the addition of methyl formate or potassium iodide enhances the amount of methanol produced. In the same table results obtained with hexacyanoferrate(II), penta- cyanonitrosylferrate(II1) and amminepentacyanoferrate(r1) complexes in place of aquapentacyanoferrate(I1) are also given. Neither hexacyanoferrate(r1) (run 8) nor pentacyanonitrosylferrate(Ir1) (run 9) is active as a homogeneous catalyst, but amminepentacyanoferrate(i1) (run 10) is. In run 11 the electrolytic solution was not bubbled with CO but contained 5 mmol dmP3 aquapentacyanoferrate(I1) and 20 mmol dm-3 methyl formate.Methanol is produced to some extent under these conditions; however, its formation was not observed in solutions without penta- cyanoferrate(I1). In all runs the amounts of formaldehyde and formic acid produced were negligibly small. Tables 1 and 2 show that an effective complex for reaction (2) must have a labile ligand to cause a substitution reaction. We therefore consider that insertion of CO occurs between Fe and the ligand, or a CO-methanol species is substituted for the labile ligand. The action of CH,OH and Fe(CN):- can be explained by a mechanism similar to that15 proposed for the catalytic conversion of methanol into ethanol in the methyl homologation reaction.16, l7 Scheme 1 shows details of the net reaction (2). Added CH,OH and Fe(CN):- operate as homogeneous catalysts, and the formation of methyl formate (believed to co cHp" I CN \ 3- kES + H+ HCHO- ES+ H+Y Scheme 1.K.OGURA AND S. YAMASAKI 27 1 be produced in the activation of the methyl group in the homologation reaction) is assumed.ls Support for this assumption is based on the experimental results that the addition of methyl formate enhanced the formation of methanol, and moreover that methyl formate itself was convertible into methanol from solution with no CO gas in the presence but not in the absence of pentacyanoferrate(I1). In CO conversion under open-circuit conditions methanol formation terminates when the oxidation of ES reaches equilibrium. However, conversion continues with the consumption of a proton in solution under controlled potential conditions.Note that this external energy is consumed not by the direct electrochemical reduction of CO, which takes place with difficulty, but by the reduction of PB to ES. ' Chem. Eng. News, 1984, 62, 8. E. Andibert and A. Raineau, Ind. Eng. Chem., 1928, 20, 1 105. H. H. Sorch, H. Golumbic and R. B. Anderson, The Fischer-Tropsh and Related Syntheses (Wiley, New York, 1951). G. Silvestri, S. Gambino, G. Filardo, M. Gauinazzi and R. Ercoli, Gazz. Chim. Ital., 1972, 102, 818. G. Silvestri, S. Gambino, G. Filardo, G. Sparado and L. Palmisano, Electrochim. Acta, 1978,23,413. fi G. Bockmair and H. P. Z. Fritz, Z. Naturforsch., Teil B, 1975, 30, 330. '' G. N. Petrova, 0. N. Efimov and V. V. Strelets, Izv. Akad. Nauk SSSR, Ser. Khim., 1983, 2042. j3 G. A. Kolyagin, V. G. Danilov, V. L. Kornienko, I. A. Kedrinskii and S . P. Gubin, Elektrokhimiya, 1983, 19, 1004. F. A. Uribe, P. R. Sharp and A. J. Bard, J. Electroanal. Chem., 1983, 152, 173. D. Ellis, M. Eckhoff and V. D. Neff, J. Phys. Chem., 1981, 85, 1225. lo V. D. Neff, J. Electrochem. Soc., 1978, 125, 886. l2 K. Itaya, H. Akahoshi and S. Toshima, J. Electrochem. Soc., 1982, 129, 1498. l 3 W. Hieber, R. Nast and C. Barstenstein, 2. Anorg. Chem., 1953, 272, 32. l 4 The Colorimetric Analytical Method, ed. L. C. Thomas and G. J. Chamberlin ( l5 Chem. Eng. News, 1982,60,41. i6 M. J. Chen and H. M. Feder, in Catalysis of Organic Reactions, ed. W. R. Moser *' M. E. Fakley and R. A. Head, App. Catal., 1983, 5, 3. Salisbury, 9th edn, 1980). New York, 1981), p. 273. M. J. Chen, H. M. Feder and J. W. Rathke, J. Am. Chem. SOC., 1982,104, 7346. rintometer Ltd, Marcel Dekker, (PAPER 4/992)

ISSN:0300-9599    

DOI:10.1039/F19858100267

出版商:RSC    

年代:1985

数据来源: RSC