摘要:

J. Chem. Soc., Faraday Trans. I, 1985, 81, 2095-2101 Kinetic Investigation of the Oxidation of Bromide Ions by Cobalt (111) Part 1.-The Influence of Pyridine in Acetic Acid Solvent BY JULIA SCHLOSSEROVA, MILAN HRONEC* AND VACLAV VESEL~ Faculty of Chemistry, Slovak Technical University, 8 12 37 Bratislava, Czechoslovakia Received 26th October, 1984 A kinetic investigation of the reduction of cobalt(rr1) acetate by bromide ions in acetic acid and in the presence of pyridine at 60-93 "C has shown that the initial rate obeys the rate law A mechanism is proposed involving the formation of bromine anion radicals as intermediates. Pyridine or other amines influence mainly (i) the actual concentration of bromide ions, (ii) the dissociation of active dimeric ColI1 species and (iii) the exchange of bromide ligands.Many industrial processes use the accelerating effect of the bromide ion on the rate of oxidation of aromatic hydrocarbons catalysed by cobalt salts in carboxylic acid media. The active catalyst is considered to be acetatobromocobalt(I1) (AcOCo'IBr).l When cobalt(I1) is oxidized to cobalt(IIr), e.g. by reaction with peroxides, the resulting complex is very reactive and abstracts hydrogen from hydrocarbons : AcOCoII'Br + RH --+ AcOCoII + R' + HBr or else decomposes :2 AcOCoIIIBr -+ AcOCoTT + Br' The bromine atoms so formed can either abstract hydrogen or recombine: Br'+RH + R+HBr Br' + Br' -+ Br,. The problem is to distinguish between abstraction of hydrogen by bromine radicals and by the cobalt bromide complex. However, in both reactions the common step is the oxidation of a bromide ion by a ColI1 ion.The kinetics and mechanism of this step have been studied in aqueous media3-5 and in acetic In aqueous media, where the cobalt ion is aquated, the oxidation of Br- occurs through the formation of [CoIII-Br-] complexes in rapid pre-equilibria followed by slow oxidation steps, and the radical ions Br; are important intermediates in this reaction. In acetic acid the reaction of CoIII ions with chloride ions proceeds through an intermediate hexachlorocobaltate(rr1) complex6? and with bromide ions through a two-stage process.8 In the latter case the first stage of the reaction is very rapid and is suggested to be the reaction of Br- with the ionic complex [CO(ACOH),]~+-(ACO-),. 20952096 OXIDATION OF Br- BY ColI1 The second, slower stage is first order with respect to both ColI1 and Br- concentration and is determined by the exchange of AcO- and Br- ligands. Amines have a pronounced effect on the activity of cobalt-bromide catalysts in acetic acid.9 They participate in the formation of cobalt-bromine complexes and probably also in the direct reaction with bromide atoms.In this paper we describe the kinetics and mechanism of the oxidation of bromide ions by cobalt(rr1) acetate in acetic acid in the presence of pyridine and some other amines. EXPERIMENTAL MATERIALS Acetic acid was purified by distillation after being heated under reflux with anhydrous chromic acid. Cobalt(I1) acetate was of reagent grade. Cobalt(iI1) acetate was prepared by the addition of an equimolar quantity of peracetic acid to a suspension of cobalt(I1) acetate in acetic acid.1° The solution was evaporated at 35-40 "C and the solid obtained dried in a vacuum dessicator over P,O,.The conversion to CoI'I was determined by cerimetric titration. The acetic acid solutions of cobalt(Ir1) acetate prepared in this manner were used for the reaction. Molecular bromine and a 47% aqueous solution of HBr were of analytical-grade purity. Pyridine and other amines were purified by distillation or crystallization. PROCEDURE The reaction was carried out in a cylindrical glass vessel placed in a thermostat. The solution, total volume of 100 cm3, was stirred at a rate of 3000 r.p.m. After the solution of hydrogen bromide, cobalt(I1) acetate, pyridine or other amines and reagents had reached the reaction temperature, a concentrated solution of cobalt(n1) acetate was added and samples of the reaction mixture were withdrawn at appropriate intervals.All experiments and analyses were carried out in nitrogen atmosphere free of oxygen. Initial rates of Co'I' and Br- consumption were obtained from numerically treated conversion curves at conversions < 10%. ANALYSIS The concentration of cobalt(II1) acetate was determined by cerimetric titration.*I l1 The concentration of bromide ions was analysed by potentiometric titration with a silver electrode coated with silver bromide using a 0.05 mol dmP3 silver nitrate solution. RESULTS KINETICS Kinetic studies were carried out under purified nitrogen in anhydrous acetic acid, mostly at 80 "C.At this temperature the rate of reduction of cobalt(rI1) acetate is very low and not influenced by the presence of pyridine or molecular bromine (see table 1). -4ddition of hydrogen bromide to an acetic acid solution of cobalt(II1) acetate resulted in the rapid reduction of ColI1 and the formation of molecular bromine. As will be shown below, when pyridine or other amines are present the initial reduction rates are decreased. The initial rate of reduction of cobalt(Ir1) acetate by bromide anion and the initial rate of Br- consumption were determined for between 6 and 9 initial concentrations of cobalt(r1) and cobalt(rI1) acetate, hydrogen bromide and pyridine. The amount of water present in the used hydrogen bromide has only a small influence (3-5%) on the rate of reduction of CoI1*. The apparent orders of reaction with respect to each component were calculated from the initial rates of consumption of CoIII and Br- ions in different concentration ranges (see table 2).The values obtained are close toJ. SCHLOSSEROV,~, M. HRONEC AND v. VESELY 2097 Table 1. Initial rate of reduction of Co'I' (ro) in acetic acid at 80 "C in the presence of different substratesa system molar ratio ro/ lo4 mol dmW3 s-l C0"I - CO"' + Py 1:4 CoIII + HBrb 1:2 Co'I' + Br, 1: 1.13 CoII' + HBrb + Py 1:2:2 0.01 0.07 23.70 0.06 3.50 a [CoII'] = 2.78 x mol dm-3, [ C O ~ ~ ] = 1.67 x lo-, mol dm-3. In the form of 47% aqueous HBr. Table 2. Apparent reaction orders for the rate of reduction of CoI'I (nco) and disappearance of Br- (nBr) reactant reaction orders concentration range / mol dm-3 nco nBr co balt(m) acetate co balt(I1) acetate HBra PY 1.39-5.56 + 1.89k0.17 + 2.10 f 0.09 1.67-8.34 - 1.lSf0.19 - 0.98 & 0.08 1.39-8.34 + 1.79k0.14 + 1.81 k0.12 1.39-1 0.42 -0.30 & 0.1 1 - 0.50 f 0.09 10.42-5 5.60 - 2.13 k 0.17 - 2.05 +_ 0.07 a In the form of 47% aqueous HBr.1 and 2, respectively, and the rate equations for higher concentrations of pyridine are as follows: (ii) d[Br-] [CoI1II2 [HBrI2 The reduction ofcobalt(II1) acetate with bromide anions usually proceeds to 70-80% conversion and pyridine and cobalt(I1) acetate were shown to have a retarding effect on the reaction rate. As illustrated in fig. 1 and 2, the initial rates vary linearly with the concentration. The calculated initial rate constants at 80 "C are k , = 4.5 x s-l and k, = 4.7 x s-l.The Arrhenius plot for oxidation of Br- and reduction of ColIJ is shown in fig. 3. The apparent activation enthalpies obtained in the temperature interval 60-93 "C are 85.3 f 3.1 kJ mol-1 for the reduction of cobalt(Ir1) acetate and 79.9 The decrease of the rate of reduction of CoIII in the presence of pyridine is also observed with other nitrogen-containing compounds (see table 3). Their action is strongly dependent on the type of nitrogen-containing compound. The ratio of the -Wo = kb [CoI*][Py]2 * 4.0 kJ mol-l for the oxidation of Br-. 69 FAR 812098 OXIDATION OF Br- BY CoIII o/ 0 0 0 / 0 o o o 0 0 0 0 0.01 0.02 Fig. 1. Plot of - (d[CoIII]/dt), against ~ , [ C O ~ ~ ~ ] ~ [HBr]z/[Coll] [PyI2.0 c 0.01 0.02 Fig. 2. Plot of - (d[Br-]/dt), against ~,,[CO~~']~ [HBrI2/[Co1I] [PyI2.J. SCHLOSSEROVA, M. HRONEC AND v. VESEL~ 2099 Y c -3.0 - - 2 . 0 2.7 2.8 2.9 3.0 lo3 KIT Fig. 3. Effect of temperature on the rate of consumption of Co'II (0) and bromide anions (A). average amount of consumed ColI1 and Br- ions (ACoIIIIABr-) remains ca. 1.07-1.10 for all amines over most of the reaction, indicating that the overall reaction closely follows CoIII + Br- + CoII +iBr,. The product of the reaction of CoIII ions with Br- is molecular bromine, which at 80 "C in acetic acid does not react with CoI1* (see table 1). Table 3. Effect of amines on the initial rate of reduction of CO"' in acetic acid at 80 "CU amine ro/ lo4 mol dmW3 s-l amine ro/ 104 mol dm-3 s-l isoquinoline 2,6-lutidine quinoline 2,4-lutidine pyridine N-oxide 2-picoline 3-picoline picolinic acid morpholine 0.06 0.09 0.13 0.15 0.40 0.45 0.50 1.01 2.13 eth ylenediamine pyridine ethanolamine 2-brompyridine nicotineamide acridine aniline without amine 2.40 5.00 6.10 7.80 9.90 9.90 < 0.01 23.70 [CO'~~] = 2.78 x mol dm-3, [CO'~] = 1.67 x mol dm-3, [HBr] = 5.56 x lop2 mol dm-3, [amine] = 5.56 x mol dm-3; nitrogen atmosphere.69-22100 OXIDATION OF Br- BY ColI1 DISCUSSION REACTION MECHANISM In order to explain the kinetic results for the initial stage of the reaction the following simplified reaction scheme is suggested : k2 CoDII1 Br + 2Py -+ CoDIIr Py + PyHBr (2) K3 CoDIIr Br + Br- $ ColI1 Br; + CoII (3) k4 CoD'I1 Br; -+ CO" + Br, where CoDII1 denotes a dimer.On the basis of literature data,2v12 the active cobalt(Ir1) species are considered to be dimers. The mechanism involves complexes formed in rapid pre-equilibria, although their structure is hard to define. When the rate-determining step is reaction (4), the initial rate of reduction of CoIII and the initial rate of Br- consumption is: -- d[ColI1] d[Br-] k , K3 k, [ C O ~ ~ I ] ~ [Br-I2 ( dt lo = - ( 7 1 , = 7 [CoII] (k; + k,[Py12) a (iii) Since we were not able to detect the proposed intermediate bromocomplexes or Br; by the e.s.r. spin-trapping technique, our insight into this mechanism has to be provided by additional experimental observations. Hydrogen bromide added to the system in the presence of pyridine or other amines is almost completely converted, e.g.to pyridinium hydrobromide : HBr + Py $ PyHBr K , PyHbr + PyH+ + Br- PyH+ + OAc- t Py + AcOH (8) which in solution is dissociated into ions, i.e. the actual concentration of Br- in acetic acid is dependent on the equilibrium constant K7 for the corresponding amine. At higher amine concentrations relative to HBr, the amine can form complexes with CoIII species [reaction (2)] and so change their activity. It means that pyridine or other amines can influence both the concentration of bromide anions and the activity of cobalt(m) acetate. This is shown by the dependence of the reaction order for pyridine in different concentration regions (see table 2). Thus, at low pyridine concentrations, reaction (6) proceeds and in eqn (iii) the reverse reaction rate constant k; is larger than k,[Py12, i.e.the reaction order with respect to pyridine will tend to zero, as was observed experimentally. It is very surprising that the presence of CoII ions retards the reduction of cobalt(II1) acetate. The retarding effect of CoI1 on the activity of cobalt(II1) acetate in acetic acid was studied during the oxidation of hydrocarbons and was ascribed to the formation of a mixed dinuclear CoI1-Co1I1 species, which are inactive for the direct abstractionJ. SCHLOSSEROVA, M. HRONEC AND V. VESELY 2101 of hydrogen from hydrocarbons.12 This assumption is based on the fact that during the oxidation of hydrocarbons, the concentration of ColI1 does not decrease below a level corresponding almost to a Corrl : CoII ratio of 1 : 1. However, in the investigated system the reduction of CoIII by bromide proceeds to 70-80% conversion and therefore the formation of less active mixed CoIII-CoII complexes does not play a significant role.The inverse dependence of the reaction rate on Corl ions can be accounted for by assuming that the reaction of a dinuclear ColI1 complex with Br- yields an intermediate mononuclear species [reaction (3)] containing a bromine anion radical in the coordination sphere. These species stabilize the Br; anion radical, which can reoxidize the Corl ion in analogy to the reaction of transition-metal ions in oxidation state(I1) by halogen radical 1 3 7 l4 Thus, in aqueous media reactions between the halogen radical anions Cl;, Br; and I; generated by flash photolysis and metal@) ions were observed, but there was no evidence of reaction between Br; and CoII: Br; + Cok; Co&I + 2Br-. (9) This has been explained14 by the unfavourable equilibrium constant for reaction (9) and the smaller standard reduction potential of Br;.However, as usual the potentials and equilibrium constants are sensitive to changes in the complexing medium.15 Therefore, in acetic acid solvent these values will change. Moreover, amine present in the solvent forms more stable complexes with ColI1 ions with lower redox potentials.16 This supports the experimental fact that with an increasing concen- tration of pyridine the reaction rate, and also the conversion of cobalt(1Ir) acetate, decreases. We suggest that in the presence of pyridine or other amines (i) exchange of bromide ligands proceeds [reaction (2)], (ii) dissociation of active dimeric CoIII species is enhanced [reaction (5)] and (iii) the positions for coordination of bromide ions with CoIII are occupied by pyridine.The kinetic data obtained in acetic acid solvent in the presence of pyridine differ from previous results obtained in the absence of pyridine* or in aqueous In both systems the reaction rate is proportional to both ColIr and Br- concentrations. In the latter case the formation of bromine anion radicals is assumed. A. S. Hay and H. S. Blanchard, Can. J. Chem., 1965, 43, 1305. R. A. Sheldon and J. K. Kochi, Metal Catalyzed Oxidation of Organic Compounds (Academic Press, New York, 1981). C. F. Wells, A. F. M. Nazer and D. Mays, J . Znorg. Nucl. Chem. 1977, 39, 2001. N . N. Malik, J. Hill and A. McAuley, J . Chem. Soc. A, 1970, 643. C. F. Wells and D. Mays, J. Chem. SOC. A, 1968, 577. B. Banas, M. Wrotiska and J. J. Ziolkowski, Roc. Chem. Ann. SOC. Chim. Pol., 1975,49, 1865. ’ A. W. Chester, E. A. Heiba, R. M. Dessau and W. J. Koehl, Inorg. Nucl. Chem. Lett., 1969, 5, 277. a V. M. Sapunov, L. Abdennur, D. Ch. Modi, 0. P. Samoplova and N. N. Lebedev, Kinet. Katal., 1975, 16, 1323. M. Hronec, Collect. Czech. Chem. Commun., 1980, 45, 1555. lo E. Koubeck and J. 0. Edwards, J . Znorg. Nucl. Chem., 1963, 25, 1401. K. Sakota, Y . Kamiya and N. Ohta, Can. J . Chem., 1969, 47, 387. l2 J. Hanotier and M. Hanotier-Bridoux, J . Mol. Catal., 1981, 12, 133. l 3 A. T. Thornton and G. S. Laurence, J . Chem. Soc., Dalton Trans., 1973, 804. l4 A. T. Thornton and G. S. Laurence, J. Chem. SOC., Dalton Trans., 1973, 1632. l5 N. Tanaka, Kogyo Kagaku Zashi, 1970,73, 2231. S. Fallab, Angew. Chem., 1967, 79, 500. (PAPER 4/ 1838)

ISSN:0300-9599    

DOI:10.1039/F19858102095

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. 1, 1985,81, 2103-2106 A Study of the Flexibility of Sodium Taurodeoxycholate Secondary Micelles in 0.8 mol d m 3 Sodium Chloride Solution BY D&AN EM~N GUVEL~~ Department of Chemistry, University of Karadeniz, Trabzon, Turkey Received 29th October, 1984 Viscosity data for sodium taurodeoxycholate, STDOC, obtained from light-scattering results (P. Schurtenberger, N. Mazer and K. Kanzig, J. Phys. Chem., 1983, 87, 303) show that the transition from primary to secondary micelles occurs at ca. 40 g dm-3 STDOC and that secondary micelles are in a flexible rod-like state in 0.8 mol dm-3 NaCl solution. It is now well established that bile salts form primary micelles in terms of the back-to-back hydrophobic bonding1-* and that larger aggregates are made of primary mi~elles.~-~ Despite these numerous reports there has been no study of the rigidity or flexibility of rod-like secondary micelles formed from bile salts in solution.For this reason we have examined the viscosity data for sodium taurodeoxycholate (STDOC) evaluated from light-scattering result^,^ and in this paper we provide new information concerning the flexibility of rod-like secondary STDOC micelles in 0.8 mol dm-3 NaCl solution. RESULTS AND DISCUSSION Since at high concentrations of bile salt and NaCl electrostatic interactions are effectively repulsive excluded-volume interactions are considered to be the major effect dominating interactions between the secondary micelles. Note that primary-secondary micelle equilibrium is not affected by the lowest-order excluded- volume interactions between rod-like micelles.' The increase in apparent mean aggregation number, nap,, at low concentrations indicates micellar growth.However, with increasing concentration naaqP reaches a maximum and then decreases at higher concentrations, showing the influence of excluded-volume interaction^.^ To account for the contribution of this effect we have deduced nap, as a function of STDOC concentration from the weight-average aggregation number, n,(C), in terms of the equation napp = n,(C)/[l+ 8ij(C-c.m.c.)] (1) where v is the partial specific volume of the STDOC monomer, C is its concentration and c.m.c. is the critical micelle concentration. Since STDOC micelles are mono- dispersed, owing to the fact that the viscosity evaluated assuming monodispersity is not significantly different than that calculated assuming polydispersity,1° intrinsic viscosities, [v], were calculated fromll [v]M/N = (1043) & (2) f Present address: 88 rue du Bac, F-75007 Paris, France. 21032104 SODIUM TAURODEOXY CHOLATE MICELLES Table 1.Hydrodynamic data for the micelles of STDOC in 0.8 mol dm-3 NaCl solution at 20 "C" (C-c.m.c.)* Rhb lvl 1?,1 (L/%pp) /gdm-3 /8, E~~~ /cm3 g-l /cm3 g-' K, A , A: L / A 18, 99.5 92.5 76.4 70.0 66.2 60.0 53.0 49.5 40.0 33.0 24.5 12.0 5.7 2.0 90 132 66.7 88 130 63.3 79 125 47.7 75" 124 41.1 72 121 37.3 70" 120 34.5 66" 116 29.9 60 113 23.1 57" 108 20.7 52" 101 16.8 45 89 12.4 36 67 8.4 30 49 6.7 21 31 3.6 40.2 38.2 28.8 24.8 22.5 20.8 18.0 13.9 12.5 10.1 7.5 5.1 4.0 2.2 52.94 24.28 23.00 749.10 5.68 50.24 23.50 22.33 724.80 5.58 37.86 19.67 19.20 606.30 4.85 32.62 17.87 17.93 551.40 4.45 29.60 16.90 16.93 518.10 4.28 27.38 15.94 16.27 492.60 4.11 23.73 14.50 15.00 447.90 3.86 18.33 16.43 13.33 9.84 6.67 5.32 2.86 a Rh is the apparent mean hydrodynamic radius; napp is the apparent mean aggregation number ; [q], the intrinsic viscosity, includes correction for the excluded-volume interactions between spherical micelles; [vC] takes into account in addition the correction for the hydration ; K,, the shape factor, includes the corrections for the excluded-volume interaction and hydration; A, is the axial ratio of a rod-like micelle, note A: is axial ratio of a micelle evaluated from Rh in terms of the equation given14 for a prolate ellipsoid by taking the semiminor axis of micelle as 15 A; L is the length of a rigid rod; L/naPp is the molecular pitch.Ref. (7). Deduced from the model of primary-secondary micelle formation by using a polymerization constant Kp = 2.7 x lo5, an aggregation number ri = 12 and a mean hydrodynamic radius Rh = 15 A for STDOC primary micelles. where &, is the apparent mean hydrodynamic radius, which is not markedly influenced by excluded-volume effects in 0.8 mol dm-3 NaC1,7 M = napp M , (where M , is the molecular weight of the STDOC monomer) and N is Avogadro's number. [r] can be generalized in the following form by including the shape factor, K,,, v and the degree of hydration, w:l2 dw is the density of water. Since cr) for sodium deoxycholate, SDOC, is independent of the concentration of added NaCl, and since presumably the value of (;o for STDOC monomer in primary aggregates is the same as that in secondary micelles because higher aggregates are formed by the association of primary rni~elles,~-~ we have derived the corresponding shape factor for the secondary micelles using the values v = 0.76 cm3 g-l and o = 0.5 g H,O g-l STDOC', from eqn (3).These are given in table 1. To include a correction for the hydration of secondary micelles, the contribution of hydration was deducted from [q] using eqn (3). On the assumption that STDOC micelles are rigid prolates, the axial ratios, A,, at STDOC concentrations > 40 g dm-3 were evaluated f r o d 3 [rl = Kc(~+w/4v)* (3) K = 1.6+- 1 + l , A R > 15 3(1n 2AR- 1.5) In 2AR-0.5 (4)D.E. GUVEL~ 2105 c Y 0 20 40 60 80 100 Fig. 1. Intrinsic viscosity plotted as a function of STDOC concentration in 0.8 mol dm-3 NaCl solution at 20 "C. Curve I takes into account the excluded-volume effects between the spherical micelles and curve I1 includes in addition the correction for the hydration of STDOC micelles. (C-c .m .c . )/g dm -3 Fig. (B) 12 13 14 \ 1 . o 4.78 4.80 4.82 4.64 log M 2. (A) Plot of M2/[qC] against log M for STDOC in 0.8 mol dm-3 NaCl solution at 20 "C; logarithmic plot of shape factor K, against log A , for STDOC secondary micelles in 0.8 mol dm-3 NaCl solution at 20 "C.2106 SODIUM TAURODEOXYCHOLATE MICELLES and are shown in table 1. The A , values compare favourably with those obtained from the equation given for Rh of a prolate ellipsoid14 by taking the semiminor axis of the secondary micelles as a 15 A (the Stokes radius of the STDOC primary micelles), and fig.1 indicates that the concentration range for the transition from primary to secondary micelles lies around 40 g dm-3 STDOC in 0.8 mol dmP3 NaCl solution, confirming a previous observation for SDOC in NaCl solution^.^ In order to explore the rigidity or flexibility of secondary micelles the following approach is adopted. It is assumed that secondary micelles are rigid rods having a length L and radius R. The rigid rod can be represented by a hydrodynamically equivalent prolate ellipsoid which has a major axis equal to L and e ual volume; the L thus obtained from the A , data are given in table 1. If secondary micelles are rigid rods their molecular pitch, L/A,,,, should be constant and independent of STDOC concentration.The pitch, L/aaPp, of a rigid rod is defined as the ratio of its molecular volume to its cross-sectional area; the radius in this case is taken to be 15 A. Table 1 shows that L/nap, increases with increasing aapp, displaying the presence of flexible micelles. On the other hand, if secondary rod-like micelles are sufficiently rigid, a plot of M2/[g,] against logM should be linear.15 It is apparent that there is a non-linear decrease in M2/[gc] as log M increases, confirming the flexibility of STDOC secondary micelles (fig. 2). For rigid rods K is proportional to Ak8 and to M1.8.11y l5, l6 Therefore the rigidity of the secondary micelles should be reflected in a plot of log K, against log A,.The slope of such a plot for STDOC at 60.0-99.5 g dm-3 in 0.8 mol dm-3 NaCl (fig. 2) is 1.53 confirming the non-rigidity of the secondary micelles. In conclusion, the results presented and discussed here demonstrate that the transition from primary to secondary micelles occurs at ca. 40 g dm-3 STDOC and that the secondary micelles are flexible and rod-like in 0.8 mol dmP3 NaCl solution. axial ratio is then A , = L/2b, where b is the semiminor axis of 15 x . The values of I thank Prof. Aykut Ikizler for helpful discussions. D. M. Small, S. A. Penkett and D. Chapman, Biochim. Biophys. Acta, 1969, 176, 178. R. Zana, J . Phys. Chem., 1978, 82, 2440. M. Vadnere, R. Natarajan and S. Lindenbaum, J. Phys. Chem., 1980,84, 1900. D. E. Guveli, J . Chim. Phys., 1984, 81, 393. D. M. Small, Adv. Chem. Ser., 1968, 84, 31. N . A. Mazer, M. C. Carey, R. F. Kwasnick and G. B. Benedek, Biochem., 1979, 18, 3064. R. Zana and D. E. Guveli, J . Phys. Chem., in press. D. E. Giiveli, to be published. ? P. Schurtenberger, N. Mazer and K. Kanzig, J. Phys. Chem., 1983,87, 308. lo R. Nagarajan, K. M. Shah and H. Hammond, Colloids Surf. 1982, 4, 147. l1 C. Tanford, Physical Chemistry of Macromolecules (Wiley, New York, 1973), p. 317. l2 J. P. Kratohvil and H. T. Delli-Colli, Fed. Proc., Fed. Am. SOC. Exp. Biol., 1970, 29, 1335. l 3 W. Kuhn and H. Kuhn, Helv. Chim. Acta, 1945, 28, 97. l4 B. Chu, Laser Light Scattering (Academic Press, New York, 1974), p. 212. l5 D. E. Giiveli, S. S. Davis and J. B. Kayes, J. Colloid Interface Sci., 1982, 86, 213. l6 J. H. Bradbury, in Physical Principles and Techniques of Protein Chemistry, ed. L. L. Leach (Academic Press, New York, 1985), part B. (PAPER 4/ 1843)

ISSN:0300-9599    

DOI:10.1039/F19858102103

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1985,81, 2107-21 14 Redox Properties of Various Bismuth Molybdate Phases in the Catalytic Oxidation of But-1-ene BY JEAN-MARIE HERRMA”* Equipe de Photocatalyse C.N.R.S., Dkpartement de Physicochimie des Materiaux, Ecole Centrale de Lyon, 35 avenue Guy de Collongues, B.P. 163, 69131 Ecully, France AND MARIA JOAO PIRES AND MANUEL F. PORTELA Laboratorio de Tecnologia quimica, Instituto Superior Tecnico, Avenida Rovisco Pais, I096 Lisboa, Portugal Received 3 1st October, 1984 The redox properties of four bismuth molybdate catalysts (a,/? and y phases and a mixture of /? and y phases) have been investigated in atmospheres of oxygen and but-1-ene at different temperatures using the electrical-conductivity technique. It has been shown that these solids are intrinsic semiconductors in the oxidized state and become highly conductive when in contact with but-1-ene.The only source of active oxygen is the surface lattice anions of the solid, whose anionic vacancies are replenished by oxygen from the gas phase. The qualitative classification of the relative reduction rates found by conductivity measurements is identical to that found by kinetic measurements carried out in a pulse reactor ( y > a > /?). Bismuth molybdates are used as selective industrial catalysts for the (amm) oxidation of olefins, and several reviews have been devoted to this system.l+ Although the unique r61e of lattice oxygen and not that of the adsorbed oxygen has been established for selective catalytic oxidation^,^ there are still problems about classifi- cation for the catalytic activity of the different M-Bi-0 phases, either for propene oxidation and ammoxidation or for the oxidative dehydrogenation of b ~ t e n e .~ The redox properties of the surface of bismuth molybdates and the lattice oxygen diffusivity are key parameters for the reactivity of these solids. Since redox processes are basically electronic, the electrical-conductivity technique appears to be a good method of following in situ the behaviour of the various Mo-Bi-0 phases when in contact with reducing or oxidizing atmospheres. The solids have previously been studied for the oxidative dehydrogenation and/or isomerization of but- 1 -ene,5* and so but-1-ene was chosen as the reducing agent for the various redox cycles.EXPERIMENTAL CATALYSTS The Mo-Bi catalysts were prepared by coprecipitation starting with a mixture of an aqueous solution of ammonium heptamolybdate and a nitric acid solution of bismuth By choosing the correct proportions and controlling the basic parameters such as pH, temperature, stirring, concentration etc., different phases of bismuth molybdate were obtained: Bi203 - 3Mo0, (a phase), Bi20;2 MOO, (jl phase), Bi203-Mo03 ( y phase) and a Mo-Bi sample, denoted MJ, corresponding to ca. 25 and 75% of the /? and y phases, respectively. All samples had low surface areas (ca. 1 m2 g-l) and were characterized by X-ray diffraction. Their structural evolution after oxidation-reduction sequences was followed by infrared spectroscopy. 21072108 REDOX PROPERTIES OF Mo-Bi CATALYSTS ELECTRICAL CONDUCTIVITY The electrical conductivity was measured in a static cell, specially designed for studying solid-gas electronic interactions at the surface of powdered oxides such as TiO,," Sn-Sb-0,12 Fe-Sr1-Sb-0'~ and metal catalysts deposited on semic~nductors.~~ The procedure and limitations of the technique have been discussed previou~ly.~~ -I4 RESULTS AND DISCUSSION SEMICONDUCTING BEHAVIOUR OF THE DIFFERENT PHASES The four samples, whose physical properties are given in table 1, were tested in electrical measurements involving the influence of the oxygen pressure Po, and of the temperature.Fig. 1 (a) shows that the electrical conductivity, 0, of the four samples previously heated in oxygen to 450 "C is independent of P02, at least in the range 2-1 60 Torr (1 Torr z 133.3 Pa), which shows that the samples are intrinsic semiconductors.Consequently, according to theory, their respective activation energies of conduction, deduced from fig. 1 (b), correspond to half their band-gap energy, EG, given in table 1. The a and y phases, which are known to be stable, have the highest EG, whereas the p phase, which is less stable, has a low band-gap energy (EG = 1.54 eV) if compared by interpolation with those of the a and y phases. The EG value of the MJ- 1 /74 sample is intermediate between those of the y and p phase, but closer to that of the p phase, although the amount of this phase in MJ is less. The immediate consequence of the intrinsic semiconductor character of the Mo-Bi catalysts is that the concentration of free electrons will be very low and no oxygen ionosorption can occur on oxidized forms close to their stoichiometric composition.Thus, the only possible source of reacting oxygen will be surface or subsurface lattice anions 02-. OXIDATION-REDUCTION IN BUT- 1 -ENEEOXYGEN SEQUENCES Since these Mo-Bi catalysts are known to oxidize but-1-ene to butadiene or isomerise it into cis- or trans-but-2-ene with two distinct patterns of behaviour depending on the temperature range (one pattern starting from 280 "C and the other ending at 350 0C),59 6 v 1 5 9 16, these two limit temperatures were chosen for the measurement of 0 during but- 1 -ene-oxygen sequences. BUT-1-ENE-OXYGEN RUNS AT 280 "C The influence upon 0 of the reduction at 280 "C by but-1-ene is given in fig.2, with a sharp initial increase followed by a slower one, especially for the a and y phases. Simultaneously, the solids become dark and behave as n-type semiconductors with the formation of anionic vacancies Vo2- : Vacancies are generally singly ionized in the ordinary thermal conditions of catalysis, thus explaining the increase of conductivity : (2) V02- + ( V&)+ + e-. The amplitude of the initial 0 variations, expressed in orders of magnitude by log (a/a,), are indicative of the relative initial rates of reduction of the various bismuth molybdates. Actually, log (a/a,) values are in the order: y > a z MJ > p. This order is identical to that obtained for the rates of reduction measured in a pulse reactor,15 as shown by fig. 3 (a). Similarly, the total amplitude 0 over ca.2 h varies with the mean rate of reduction rates of Mo-Bi by but-1-ene [fig (36)l in the order y > a > MJ > p.J-M. HERRMANN, M. J. PIRES AND M. F. PORTELA 2109 Table 1. Physical characteristics of Mo-Bi samples catalyst d logo formula phase Bi/Mo structure colour d log Poz EJkJ mol-1 E,/eV Bi@, * 3M00, a 2/3 scheelite pale yellow 0 130 2.67 1 - pale yellow 0 75 1.54 2/ 1 koechlinite bright yellow 0 146 3.04 Bi,O, - 2Mo0, P Bi,O, - MOO, Y 75%~+25%p MJ-1/74 - - bright yellow 0 102 2.12 -7 -8 n * I !j - 9 " I C ;s- -10 -11 W on - -12 a! -a a-a-a- -13)- I 1 I I 1.4 1 a 6 1.8 2.0 103 KIT Fig. 1. (a) Conductivity isotherms: plots of log o againstfIlogPO2) at 450 "C. (b) Arrhenius plot of log o againstf11O3/T) at Po, = 160 Torr.21 10 - 2 - 3 - 4 - 5 5 - & -6 v - % - 7 - a -9 -1 0 REDOX PROPERTIES OF Mo-Bi CATALYSTS 0 100 200 300 t/min Fig. 2.Plot of electrical conductivity (log a) against time with but-1-ene (60 Torr) and oxygen (1 60 Torr) at 280 "C. After prompt outgassing, the introduction of oxygen at the same temperature causes a sharp decrease of o, corresponding to the fast filling of surface anionic vacancies: iO,(g) + (V02-)+ + e- f 0:-. (3) An electrical steady state is reached rapidly in the case of the /? phase and the MJ sample. The slower o variations found for the a and y phases, for reduction as well as for reoxidation (fig. 2), seem to be indicative of deeper oxidation-reduction processes affecting the subsurface lattice oxygen. By comparing the variations of o in but-1-ene and in oxygen, it appears that only the /? phase is reoxidized completely in 0, at 280 "C.BUT- 1 -ENE-OXYGEN RUNS AT 350 "C The same sequences were repeated at 350 "C (fig. 4), which corresponds to the higher temperature used for catalytic tests. As with low-temperature experiments, the initial electrical conductivity variations (log o/ao) provide the relative, initial rate of reduction. These are determined on the vertical axis at time t = 0 and vary according to the nature of the Mo-Bi phases in the same order as obtained for initial rates of reduction measured at 350 "C in pulse reactor:15 y > a > MJ > /?. For the mean rates of reduction, determined for t z 100 min, electrical-conductivity estimates always parallel the kinetic data obtained in the pulse reactor, but with an inversion between the a and y phases: a > y > MJ > p.After rapid evacuation, the introduction of oxygen causes a sharp drop in the electrical conductivity and a simultaneous change in the colour of the catalysts fromJ-M. HERRMANN, M. J. PIRES AND M. F. PORTELA 2111 Fig. 3. Plot of relative electrical conductivity (loga/a,) as a function of rate of reduction determined in a pulse reactor: (a) initial variation of a as a function of initial reduction rates and (b) total variation of a within ca. 100 min as a function of the mean reduction rate. T = 280 "C. 02 1 0 100 200 300 tlmin Fig. 4. Same as in fig. 2 but at 350 "C.21 12 REDOX PROPERTIES OF Mo-Bi CATALYSTS , I 1-C4H8 ' 02 Fig. 5. Plot of log 0 for but-1-ene-oxygen sequences at 450 "C.black to yellow. This corresponds to the filling of peripheral anionic vacancies [reaction (3)]. Note that the initial slopes (d log a/dt)o are smaller for oxidation than for reduction. Their relative ratios (red/ox) are always higher than unity (4.4,l. 1,8.7 and 3.8 for a,P, y and MJ, respectively). This means that the rate of surface reduction of Mo-Bi by but-1-ene is higher than the rate of reoxidation by gaseous O,, as was found for the same solids in the presence of an ammoxidation reacting mixture of propene, ammonia and 0xygen.l' Consequently, we conclude that the Mo-Bi catalysts are in a partially reduced surface state during the reaction. The redox process, which involves the participation of lattice oxygen ions, is described mathematically by the Mars and Van Krevelen mechanism.l* OXIDATION-REDUCTION SEQUENCES AT 450 "C AND STRUCTURAL EVOLUTION under the reducing influence of but- 1 -ene according to Previous X-ray diffraction s t ~ d i e s ~ ~ j 2o have shown that pure Mo-Bi-0 phases react Bi,O, * 3MoO,(a) + Bi,O, - MoO,(y) + 2Mo0, + 2 0 Bi,O, - 2MoO,@) -+ Bi,O, - MoO,(y) + MOO, + 0 BI,O, * MoO,(y) + 2Bi + MOO, + 40.Moreover, when the P phase is reduced to metallic Bi and MOO,, subsequent reoxidation with oxygen does not restore the initial structure and some formation of the a-phase occurs. The results of the successive but-1 -ene-oxygen sequences on the a phase at 280,350 and 450 "C (fig. 2,4 and 5 ) show a fast initial redox process followed by a slower one, as mentioned in ref.(17), which has to be related to its close-packed structure. At lower temperature the /3 phase is only superficially reduced, while a phase sublayers are involved even at 280 "C. In the reoxidation cycle the initial a phase structure is not restored, at lowJ-M. HERRMANN, M. J. PIRES AND M. F. PORTELA 1000 800 600 400 1000 800 600 400 I wavenum ber/cm -' 21 13 1000 800 600 400 1000 800 600 400 wavenumberlcm-' Fig. 6. 1.r. spectra of (a) a, (b) p, (c) y and (6) MJ samples (-) before and (---) after redox cycles at 450 "C. temperature, during the analysed time, oxygen diffusion in the lattice being very slow : reversibility was observed at 450 "C [fig. 5 and 6(a)]. In P-phase reoxidation, reversibility was not achieved when sublayers were involved (450 "C), as shown in fig.6(b) where characteristic i.r. bands of the a and y phases are also visible. The results obtained with the y phase (fig. 2,4 and 5) show good agreement with previous data,15 suggesting that this phase is the most reducible and the most promptly stabilized : the reduction-reoxidation process is reversible at high temperatures [fig. 5 and 6(c)]. The multiphase MJ sample, composed by P and y phases, presents intermediate behaviour : the electrical conductivity at high temperatures is similar during reduction to that of its less reducible component (J phase). and is similar during reoxidation to that of its more easily oxidized component ( y phase) [fig. 5 and 6(d)]. CONCLUSIONS From the electrical-conductivity data for the four bismuth molybdates studied (a, P and y phases and mixed P + y phases) the following points can be deduced. (1) In oxygen, after preoxidation, the solids behave as intrinsic semiconductors with low conductivities, independent of oxygen pressure in the range 2-160 Torr, as would be21 14 REDOX PROPERTIES OF Mo-Bi CATALYSTS consistent with the stoichiometric composition and the absence of anionic vacancies in equilibrium with Pop.(2) In the presence of but-I-ene the solids become highly conductive because of the formation of singly ionized vacancies due to the oxidative dehydrogenation of C,H, into C,H, [reaction (l)]. The rates of reduction of the solids, deduced from variations of 0, vary in the same way as those obtained from pulse-reactor kinetics ( y > a > MJ > a).(3) Reoxidation of butene-treated samples by 0, at 280 and 350 "C causes the conductivity to decline towards values originally observed for oxygen-pretreated samples, consistent with reconstitution of surface lattice anions from anion vacancies caused by reduction. (4) From conductivity cycles in butene and oxygen it was observed that the initial rates of reduction are higher than the initial rates of oxidation, which shows that the surfaces of the solids are in a slightly reduced state (Mars and Van Krevelen mechanism). ( 5 ) Electrical-conductivity measurements and infrared spectroscopy gave complementary results showing that the ct and y phases undergo reversible surface structural changes in redox sequences at 450", in contrast to the phase. This work was carried out at the Institut de Recherches sur la Catalyse C.N.R.S., 69626 Villeurbanne, France, the former address of J-M.H.We thank the CIES for financial support (to M. J. P.). G. W. Keulks, L. D. Krenzke and T. M. Notermann, Adv. Catal., 1978, 27, 183. R. K. Grasseli and J. D. Burrington, Ado. Catal., 1981, 30, 133. A. Bielanski and J. Haber, Catal. Rev., 1979, 19, 1. D. Carson, G. Coudurier, M. Forissier, J. C. Vedrine, A. Laarif and F. ThCobald, J. Chem. Soc., Faraday Trans. I, 1983, 79, 1921. M. M. Oliveira, M. F. Portela, M. J. Pires and F. R. Ribeiro, Can. J . Chem. Eng., 1983, 61, 87. M. F. Portela, M. J. Pires and M. M. Oliveira, Proc. 8th Iberoamer. Symp. Catal., Huelva, Spain, 1982, p. 374. ' M. J. Pires, M. F. Portela, M. M. Oliveira, A. Saraiva and T. Miranda, Proc. 7th Iberoamer. Symp. Catal., La Plata, Argentina, 1980, p. 189. G. W. Keulks, J. Hall, C. Daniel and K. Suzuki, J. Catal., 1974, 34, 79. P. A. Batist, J. F. H. Bowens and G. C. A. Schuit, J. Catal., 1972, 25, 1. lo C. R. Adams, H. H. Voge, C. Z. Morgan and W. E. Armstrong, J . Catal., 1964, 3, 379. l1 J-M. Herrmann, J. Chim. Phys., 1974,73,474; 479. l2 J-M. Herrmann, J-L. Portefaix, M. Forissier, F. Figueras and P. Pichat, J. Chem. SOC., Faraday l3 J-M. Herrmann and B. Benaichouba, React. Kinet. Catal. Lett., 1983. 22, 209. l4 J-M. Herrmann and P. Pichat, J. Catal., 1982, 78, 425. l5 M. M. Oliveira, M. F. Portela, M. J. Pires, 5th Portug. Nat. Conf. Chem., Porto, Portugal, 1982. Trans. I, 1979, 75, 1346. M. F. Portela, M. M. Oliveira, M. J. Pires, F. M. S. Leonos and L. Ferreira, Proc. 8th Int. Congr. Catal., West Berlin, (Verlag Chemie, Weinheim, 1984), vol. 11, p. 533. l7 J. F. Brazdil, D. D. Suresh and R. K. Grasselli, J. Catal., 1980, 66, 347. P. Mars and D. W. Van Krevelen, Chem. Eng. Sci. Suppl., 1954, 3, 41. l9 M. Egashira, H. Sumie, T. Sakamoto and T. Seyama, Kogyo Kageku Kyokai-Shi, 1970, 73, 860. 2o M. F. Portela, Proc. 8th Iberoamer. Symp. Catal., Huelva, Spain, 1982, p. 315. (PAPER 4/ 1862)

ISSN:0300-9599    

DOI:10.1039/F19858102107

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. 1, 1985, 81, 21 15-2122 Intensely-scattering Phase in Surface-enhanced Raman Scattering by Cyanide on Gold Electrodes BY MERRICK R. MAHONEY AND RALPH P. COONEY* Chemistry Department, University of Newcastle, New South Wales 2308, Australia Received 5th November, 1984 Surface-enhanced Raman scattering (SERS) by cyanide on gold and silver electrodes in 0.1 mol dm-3 KOH, 0.1 mol dm-3 KCN appears to arise from metal oxide<yanide compounds and not from cyanometal complexes. The similarity of v(CN) for the two metals indicates that metal(ionHyanide interactions within the SERS compounds are relatively unimportant. Possible involvement of cyanogold(r) complexes (e.g. [Au(CN),]- } is ruled out by substantial differences between the v(CN) band for the SERS phase (ca.21 10 cm-') and for these complexes (> 2164 cm-'). The existence of analogous SERS compounds for both silver and gold is indicated by similar SERS frequencies under both controlled-potential and withdrawn (dry)- electrode conditions for the two metals. The involvement of a hydrolysed component (oxide/hydroxide) in gold/cyanide SERS is suggested by previous electrochemical studies of the gold dissolution system and by the similarity of gold/cyanide SERS to SERS from the system Ag/O.1 mol dm-3 Na,SO,, 0.1 mol dmV3 KCN, which appears to arise from a hydrolysed cyanide compound. Some indications of laser damage exist in the gold/cyanide SERS system, which is classified as an SERS type I11 (inorganic microzone). A recent survey1 of nine candidates for the intensely scattering phase in surface- enhanced Raman scattering (SERS) by cyanide on silver electrodes has favoured a stoichiometric silver oxide-cyanide compound.This selection was based on the ability of each species to account for eight specific experimental criteria.l The present investigation seeks to extend this analysis to the case of gold/cyanide SERS where previous electrochemical studies2 have revealed that dissolution processes involve passivating layers of gold(1, 111) oxide or hydroxide. The electrolyte in that electro- chemical investigation2 of cyanide dissolution of gold was 0.1 mol dm-3 KOH, 0.1 mol dm-3 KCN, which differs from the more commonly employed electrolyte (0.1 mol dmP3 Na2S04, 0.01 mol dma3 KCN) in cyanide SERS ~tudies.~ Therefore, to facilitate comparison of cyanide SERS from silver and gold electrodes, the former electrolyte was employed for both metals in this study.Previous gold/cyanide SERS studies4 have employed chloride-containing electrolytes. The position of the v(CN) line in gold/cyanide SERS is a key factor in the assignment of the spectrum. Because the silver SERS v(CN) line5-' (ca. 21 10 cm-l) is within the range of frequencies (table 1) of cyanosilver(1) complexes (2080-2 150 cm-1),8-11 the SERS spectrum has frequently been attributed to such complexes. However, the gold SERS v(CN) line (ca. 21 10 cm-l), which is similar to that for silver SERS, differs significantly from the range of frequencies of cyano- gold(1, 111) complexes (2 2164 cm-l, see table l).8-11 While [Au(CN),]- has been invoked as the Au/CN- SERS specie^,^ an alternative explanation is considered in this study.211521 16 SERS BY CYANIDE ON GOLD ELECTRODES Table 1. Cyano stretching-mode frequencies v(CN)/cm-l ref. 2094-21 14a ] this work 2095-2111a 2095-21 16" this work 2209 2239 2164 8-1 1 a Position varies with potential. EXPERIMENTAL The lasers, Raman spectrometer, electrochemical equipment and cells have been described previously.12 Unless otherwise indicated the electrolyte used in this investigation was aqueous 0.1 mol dm-3 KOH, 0.1 mol dm-3 KCN. Raman spectra were recorded using the 647.1 nm line of a krypton-ion laser at a power of ca. 100 mW. The electrode surface was aligned by peaking reflective scattering from the surface prior to the oxidation-reduction cycle (0.r.c.). All potentials quoted were measured relative to a saturated calomel reference (SCE).The polycrystalline gold or silver electrode (Johnson-Matthey) was polished with a fine emery paper and then etched in hot aqua regiu for 10 s. The electrode was then washed with distilled and fractionated water and placed in a spectroelectrochemical cell (following the design of Pettinger et all3) containing 0.1 mol dmP3 KOH and 0.1 mol dm-3 KCN and cathodically cleaned at - 1.2 V for at least 30 min. At the end of this period a Raman spectrum recorded from the gold or silver electrode at - 1.2 V over the region 1900-2200 cm-l revealed only a very weak band at ca. 2081 cm-' (i.e. electrolyte solution cyanide).*-'' The electrode was then subjected to an 0.r.c.which consisted of either sweeping the potential from - 1.2 to +0.55 V (usually) and back to - 1.2 V at a sweep rate of 1 mV s-l or a potential step to +0.5 V with a pause of 5 s and then back to - 1.2 V. The former are the conditions employed in recording the cyclic voltammogram in fig. 1. After such an 0.r.c. an intense SERS v(CN) line was observed at ca. 2100 cm-l. RESULTS AND DISCUSSION The dependence of the position and intensity of the v(CN) (ca. 2100 cm-l) SERS line on the anodic limit of the 0.r.c. is summarized in table 2. Several important points emerge from the data. First, an 0.r.c. with an anodic limit of -0.65 V is not sufficient to generate the intense SERS line at ca. 2100 cm-l. As -0.65 V represents the first dissolution stage2 of the gold electrode, to form [Au(CN),]-, it would appear that the formation of [Au(CN),]- per se on an anodized surface is not a sole prerequisite for the observation of SERS.(Previous studies have assigned the SERS spectrum to this specie^.^) Secondly, the observation of an intense SERS line at ca. 2100 cm-l for anM. R. MAHONEY AND R. P. COONEY 21 17 V/V us SCE J Fig. 1. Cyclic voltammogram from a polycrystalline gold electrode in aqueous 0.1 mol dm-3 KOH, 0.1 mol dm-3 KCN. Table 2. Effect of the anodic limit of the o.r.c.a on SERS intensities at ca. 2100 cm-l for the system Au/O.l mol dmT3 KOH, 0.1 mol dm-3 KCN peak intensityC anodicb at ca. 2100 cm-l anodic pause at charge limit anodic transfer 1 min after 10 min after /V us SCE limit/s /mC cm-2 0.r.c. 0.r.c.d - - - - 0.65 5-60 -0.45 5 + 0.05 5 10 14 8 +0.12 5 25 93 43 + 0.30 5 55 320 300 +0.50 5 10 74 62 - - a 0.r.c.: - 1.2 Veanodic limit; focussed 647.1 nm Kr+ was present during the 0.r.c. and Assuming roughness factor the surface was cathodically cleaned before each successive 0.r.c. of 1. Relative values. For 60 s pause, charge transfer estimated to be 51 mC cm-2. 0.r.c. anodic limit of + 0.05 V (table 2) suggests that a gold(1Ix) species (such as Au,O, formed at ca. + 0.45 V)2 is also not directly involved in the intensely scattering phase. The 0.r.c. limit (+0.05 V) for the onset of SERS is in the region of the cyclic voltammogram that has been shown to involve a one-electron-transfer process.2 Au/O.l mol dm-3 KOH, 0.1 mol dm-, KCN Prior to an oxidation-reduction cycle the Raman spectrum (fig.2) from the gold electrode surface incorporates a weak band at ca. 2081 cm-l and more intense lines at 1030 and 1064 cm-l. The former features arises from solution cyanide8-l1 while the latter two features have no counterparts in the spectrum of the electrolyte and are therefore attributed to surface species. The Raman spectrum of uncomplexed carbonate14 ion exhibits an intense v1 band (D3h:A;) at 1065 cm-l and complexed carbonate15 exhibits a related mode at lower frequencies. Therefore the 1030 and 1064 cm-l features are attributed to surface carbonates. After an o.r.c., two additional lines appear (fig. 2) at ca. 2100 and 530 cm-l, which ha.ve no counterparts in the spectrum of the bulk electrolyte. The weak, broad21 18 SERS BY CYANIDE ON GOLD ELECTRODES 200 1200 2200 3200 Raman shift/cm-’ Fig.2. Raman spectra from a polycrystalline gold electrode in aqueous 0.1 mol dmP3 KOH, 0.1 mol dm-3 KCN. The sharp line near 2850 cm-l is an instrumental artifact. (a) - 1.2 V before an o.r.c., (6) +0.45 V after an 0.r.c. and (c) - 1.2 V after an 0.r.c. underlying profile of reduced carbon (1 100-1 700 cm-l) observedlG in type I SERS (e.g. Ag/organic adsorbate) is not evident in these spectra. However, intensified carbonate lines together with a persistent fluorescence continuum were evident (fig. 2) at + 0.45 V after an 0.r.c. A similar continuum is characteristic of highly oxidized carbon (graphite oxide/acid)” which may form in the present system from carbonate decomposition. It would appear that the oxidized carbon is not a component of the intensely scattering SERS phase because at more cathodic potentials (- 1.2 V) following an o.r.c., v(CN) (ca.2100 cm-l) has gained in intensity and both carbonate (1030 and 1064 cm-l) and the continuum have diminished in intensity (fig. 2). Also, an intense 2100 cm-l line has been observed after a potential-step 0.r.c. in a system which exhibits no apparent evidence for carbonates. Because the 530 cm-l line (fig. 2) does not follow a constant intensity relative to the 2100 cm-l line, it is not considered to be a different mode of a common scattering species. The 530 cm-l line has only been observed under electrochemical conditions known to generate a surface film of gold(I1r) oxide on a gold electrode surface2 in the absence of laser flux.The infrared spectrum of gold(m) oxidela exhibits absorptions in the region 660-517 cm-l and therefore the 530 cm-1 line is attributed tentatively to an internal mode of gold(rr1) oxide or to a related hydrolysed phase. Evidence will be presented below which suggests that the ca. 2 100 cm-1 line arises from a hydrolysed compound of gold(1) cyanide. Both the intensity and frequency of the ca. 2100 cm-l line are sensitive to the applied potential (table 3 and fig. 3). The variation in SERS intensity with changing potential, which has a maximum at ca. - 1.1 V, was complicated by the rise and fall of band intensity with time (fig. 4). The variation in frequency with changing potential (table 3) parallels that observed for Ag/cyanide SERS, 1 which suggests a common scattering species for cyanide on both metals.After a double potential-step 0.r.c. intensity against time curves for a gold surfaceM. R. MAHONEY AND R. P. COONEY 21 19 Table 3. Dependence of SERS frequency and line position on potential after the 0.r.c. for Au/O.l mol dm-3 KOH, 0.1 mol dmP3 KCN potential v(CN) relative /v /cm-l peak heighta -1.30 2095 85 -1.20 2100 92 -1.00 2107 100 -0.80 2116 20 b - - 3 -0.65 a Approximate intensities (arbitrary units). Not observed. L L L I I 1 I I 1 I 1 2000 2100 2 200 Raman shift/cm-' Fig. 3. Raman spectra from the surface of a polycrystalline gold electrode in aqueous 0.1 mol dm-3 KOH, 0.1 mol dm-3 KCN. All spectra were recorded after an 0.r.c. : (a) - 0.8, (b) - 1.0, ( c ) - 1.2 and ( d ) - 1.3 V.subjected to laser interruption provided no clear evidence for laser-assisted dissolution processes found for Ag/cyanide SERS. However, uniform low laser flux (at 5 1 % of SERS flux)6 induced perturbations of cell currents. At conventional SERS flux levels (100 mW of focused 647.1 nm Kr+) the perturbations within the laser sampling zone induced by the laser flux would be expected to be increased substantially. This result is similar to that previously reported for Ag/CN- SERS.6 As has been found for Ag/cyanide SERS,5f l6 the spectrum from the dry (withdrawn) gold electrode (fig. 5) differed in significant respects from the spectrum recorded from the gold electrode under working cell conditions. It was also apparent that the dry-electrode spectra varied depending on the 0.r.c.employed. After a potential sweep 0.r.c. to +0.50 V, the Raman spectrum of the dry gold electrode incorporated an intense v(CN) band at ca. 2138 cm-l together with the carbonate bands at 1025 and 1063 cm-l. In contrast to the spectrum from the withdrawn electrode exhibiting Ag/cyanide SERS,5y l6 no carbon bands (1 100-1 700 cm-l) were evident in Au/cyanide SERS. Several unassigned low-frequency bands were also observed at 389, 478 and 578 cm-l. The 2138 cm-l line in the dry gold electrode spectrum is displaced 58 cm-l from the free cyanide ion frequency. This displacement [Av(CN)] is similar to that for the silver/cyanide SERS system [Av(CN) x 64 cm-l]. The values for silver [Av(CN) x 30 cm-'1 and gold [Av(CN) z 25 cm-'1 SERS under working electrode2120 SERS BY CYANIDE ON GOLD ELECTRODES 200 x c c c m ._ * .d .z 100 Y m - 2 0 ' 60 120 r/min Fig.4. Variation of the intensity of the line at ca. 2100 cm-l as a function of time after the 0.r.c. for a polycrystalline gold electrode in aqueous 0.1 mol dm-3 KOH, 0.1 mol dm-3 KCN at - 1.2 V. I I I I I I I I I I I 1 200 1000 2000 2400 Raman shift/cm-* Fig. 5. Raman spectrum from a dry (withdrawn) polycrystalline gold electrode after a potential-sweep 0.r.c. in aqueous 0.1 mol dmP3 KOH, 0.1 mol dm-3 KCN.M. R. MAHONEY AND R. P. COONEY 2121 conditions are very similar and the 2138 cm-l line is similar to the dominant cyanide SERS line from gold sol surfaces.19 The spectrum of the dry-electrode surface after a potential sweep to +0.3 V did not exhibit a v(CN) SERS line.Because the step to +0.55 V produces2 an oxide passivating film on the electrode surface (in the absence of laser flux) and a step to + 0.3 V does not,, it appears probable that the appearance of gold/cyanide SERS from a dry electrode is predicated on, but not directly associated with, the formation of a gold(II1) oxide phase. It is possible that the gold(1) oxide-cyanide compound (see below) forms from the gold(II1) oxide precursor. Ag/O. 1 mol dm-3 KOH, 0.1 mol dm-3 KCN To facilitate comparison of SERS from cyanide on silver and gold electrodes, spectra were also recorded from silver electrodes in the electrolyte 0.1 mol dm-3 KOH, 0.1 mol dm-3 KCN. Similar 0.r.c. pretreatments were used for both metals. After polishing, etching and cathodic cleaning (- 1.4 V for ca.60 min) the spectrum from the silver surface exhibited no features attributable to surface species. After an 0.r.c. intense v(CN) band was observed at 2095-21 11 cm-l with the precise position depending on potentia1.l This is very similar to the range observed' for the system Ag/O. 1 mol dm-3 Na,SO,, 0.0 1 mol dmP3 KCN. The similar range suggests that a common scattering species exists within the laser microzone of the silver surfaces for both electrolytes.' The cathodic displacement of the potential of maximum SERS intensity on changing from a Na,SO,- to a KOH-containing electrolyte is consistent with the stabilization of an oxide/hydroxide scattering phase at the higher pH of the latter e1ectrolyte.l' Furthermore, the very similar v(CN) range (2095-2 116 cm-l) for both silver and gold electrodes in the KOH-containing electrolytes strongly suggests, given the assignment data in table 1, that the SERS phases for both silver and gold are closely related.The insensitivity of the v(CN) band to the change from silver to gold, which would be expected to generate a pronounced frequency shift for direct metal-cyanide interactions (see table 1) indicates that such interactions are relatively unimportant in the SERS phase. Some differences were observed between the silver and gold systems that do not appear to negate the conclusions above. The cyclic voltammograms recorded during the: 0.r.c. [ - 1.2 V t + 0.55 V (at 1 mV s-l)] revealed that, at potentials more anodic than -0.5 V, there is a current plateau ( i e .there is no evidence for passivation of the type observed for the gold system). This difference in passivation is likely to originate in different solubilities for the related oxide-cyanide compounds suggested by the data discussed earlier. Also, the SERS spectra for the Ag/O. 1 mol dm-3 KOH, 0.1 rnol dm-3 KCN were obtained with comparable intensity and constant v(CN) using both 514.5 nm Ar+ and 647.1 nm Kr+. This is in contrast with the results for gold electrodes in the same media which reveal that SERS is observed with 647.1 nm Kr+ but not with 514.5 nm Ar+. Given the extensive evidence for laser- damage processes5~ in the SERS system Ag/O.l mol dm-3 Na,SO,, 0.01 mol dm-3 KCN and the evidence for similar damage in the Au/O.l mol dm-3 KOH, 0.1 mol dm-3 KCN case, it appears probable that such excitation differences originate in differences in photon-surface resonance.lG Such resonance is thought to cause the increased surface corrosion within the laser microzone by creating a localized ' warm ' zone.5'2122 SERS BY CYANIDE ON GOLD ELECTRODES CONCLUSION THE IDENTITY OF THE GOLD/CYANIDE SERS PHASE Overall, the indications from this study are that the SERS phases in both silver/cyanide and gold/cyanide systems are closely related.Therefore in the light of the systematic analysis' of silver/cyanide SERS, it would appear that the SERS phase in both cases is an oxide-cyanide stoichiometric phase (i.e. a compound). While the study of laser-damage effects in the gold/cyanide SERS is only at a preliminary stage, it appears on the basis of the uniform (low) flux perturbations that laser damage may occur.This is not unexpected in the light of the extensive indications of laser damage in the silver/cyanide SERS We thank the Australian Research Grants Scheme for providing the equipment for this project. M. R. Mahoney and R. P. Cooney, Chem. Phys. Lett., accepted for publication. D. W. Kirk, F. R. Foulkes and W. F. Graydon, J. Electrochem. Soc., 1980, 127, 1962. R. P. Cooney, M. R. Mahoney and A. J. McQuillan, in Advances in Infraredand Raman Spectroscopy, ed. R. J. H. Clark and R. E. Hester (Heyden, London, 1982), vol. 9, pp. 188-281. H. Baltruschat and J. Heitbaum, J. Electroanal. Chem., 1983, 157, 319. R. P. Cooney, T. P. Mernagh, M. R. Mahoney and J. A. Spink, J. Phys. Chem., 1983, 87, 5314. M. R. Mahoney and R. P. Cooney, J. Phys. Chem., 1983,87,4589. M. R. Mahoney and R. P. Cooney, J. Raman Spectrosc., 1981, 11, 141. D. M. Adams, Metal-Ligand and Related Vibrations (Edward Arnold, London, 1967), pp. 164-168. K. Nakamoto, Infrared Spectra of Inorganic and Coordination Compounds (Wiley-Interscience, New York, 1978), chap. 3. L. H. Jones, J. Chem. Phys., 1965,43, 594. lo R. A. Penneman and L. H. Jones, J. Chem. Phys., 1958, 28, 169. l2 M. R. Mahoney, M. W. Howard and R. P. Cooney, Chem. Phys. Lett., 1980,71, 59. l3 B. Pettinger, U. Wenning and D. M. Kolb, Ber. Bunsenges. Phys. Chem., 1978, 82, 1326. l4 S. D. Ross, Inorganic Infrared and Raman Spectra (McGraw-Hill, New York, 1972), p. 160. l5 L. H. Little, Infrared Spectra of Adsorbed Species (Academic Press, New York, 1964). l6 R. P. Cooney and T. P. Mernagh, in Electrochemistry - The Interfacing Science, ed. A. M. Bond and D. A. J. Rand (Elsevier, Amsterdam, 1984) (also published in J. Electroanal. Chem., 1984, 168, 67). T. P. Mernagh, R. P. Cooney and J. A. Spink, J. Chem. SOC., Faraday Trans. I , 1984, 80, 3469. E. Schwarzmann and G. Gramann, 2. Naturforsch., Teil B, 1970, 25, 1308. K. U. Von Raben, R. K. Chang and B. L. Laube, Chem. Phys. Lett., 1981,79,465. (PAPER 4/ 1885)

ISSN:0300-9599    

DOI:10.1039/F19858102115

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年代:1985

数据来源: RSC

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J . Chem. SOC., Faraday Trans. I , 1985,81, 2123-2130 Chemical Origins of Surface-enhanced Raman Scattering by Cyanide on Copper Electrodes BY MERRICK R. MAHONEY AND RALPH P. COONEY* Chemistry Department, University of Newcastle, New South Wales 2308, Australia Received 5th November, 1984 A simple potential-step oxidation-reduction cycle pretreatment is reported for the surface- enhanced Raman scattering (SERS) system Cu/O.1 mol dm-3 Na,SO,, 0.01 mol dm-3 KCN. Laser-assisted corrosion and dissolution processes are suggested by a combination of uniform- laser-flux perturbations of the cyclic voltammetry and laser-interruption effects in the intensity against time decay curve for the system at - 1.25 V (us SCE). The pH dependence of v(CN) intensities for both the copper/cyanide SERS surface and for related cyanocopper(1) complexes is attributed partly to the decomposition of such complexes (pH < 4.6) and partly to the decomposition of associated copper oxides (pH > 2).The identification of cyanocopper(1) complexes stabilized by copper(1) oxide as the intensely-scattering phase compares with the oxidexyanide phase proposed for silver/cyanide SERS. The copper/cyanide SERS system is classified as type I11 (inorganic microzone). Previous studies1 of surface-enhanced Raman scattering (SERS) by cyanide on silver and gold electrodes suggest that the intensely-scattering phase is a metal oxide-cyanide compound. These and other studies suggested that the appearance of SERS in many electrochemical systems correlates with laser-induced compositional andmorphological damage ofthemetal Thecomposition of the laser-damage microzone has been employed2* as the basis of a classification of SERS, uiz. type I, (adsorbate) carbon ; type 11, (non-adsorbate) carbon ; type 111, inorganic phases ; type IV, non-enhanced and type V, resonance (non-surface) enhanced. On this basis, silver/cyanide and gold/cyanide SERS were classified as type 111.The present study of cyanide SERS from copper electrodes has as its two objectives the identification of the intensely-scattering phase and the testing for laser-induced copper surface damage. Cyanide SERS from copper electrodes has been less extensively studied6-8 than the corresponding silver system. The Raman spectrum from a dry (withdrawn)-electrode surface6 was very similar to bulk CuCN(s), which suggests that the surface is covered by multilayers of CuCN.Under controlled-potential conditions SERS was observed only when exciting wavelengths > 575 nm were employed.' The oxidation-reduction cycle pretreatment was thought to generate a relatively dense layer of copper@) complexes on the surface.' While previous studies of this system had attributed the SERS spectrum to cyanocopper(1) complexes,8 this possibility was further tested in the present study in the light of the contrasting assignment (to an oxide-cyanide compound) made for the corresponding silver system. 21232124 SERS BY CYANIDE ON COPPER ELECTRODES ~ l l l l l l l l l ~ 2000 2100 2200 Raman shift/cm-' Fig. 1. Variation of the Raman spectrum from the surface of a copper electrode with time after the 0.r.c.Electrode potential: -0.6 V; electrolyte: 0.1 mol dmP3 Na,SO,, 0.01 mol dm-3 KCN. (a) 10, (6) 295 and (c) 595 s after the 0.r.c. EXPERIMENTAL The lasers, Raman spectrometer, electrochemical equipment cell design and reagents have been described previously.3* Distilled water, refractionated under N,, was employed in preparing electrolyte solutions. Copper foil (B.D.H. AnalaR foil) and 1 mm diameter copper wire (Alfa Products, m3N5) were used as electrode materials. The 1 mm diameter wire was used in the uniform-laser-flux perturbation studies. In these uniform-flux studies the unfocused laser beam (1.4 mm diameter) filled the wire cross-section which served as the working electrode surface. Unless otherwise stated the electrolyte employed in these studies was aqueous 0.1 mol dm-3 Na,SO,, 0.01 mol dm-3 KCN.Raman spectra were generally recorded using 647.1 nm Kr+ at 100 mW output power. For all studies except the uniform-flux perturbation experiments, the laser beam was focused to a point (ca. 0.04 mm in diameter as indicated by scanning Auger microscopy* and theoretical calculationslO) on the electrode surface. The band-pass used was 10 cm-l. The oxidation-reduction cycle (0.r.c.) pretreatment is described below. Prior to an 0.r.c. the electrolyte was thoroughly deoxygenated by bubbling N, through it (ca. 20 h), and the copper working electrode was held at - 1.6 V for 3&60 min in the electrolyte to be studied. Such precathodization ensured that the electrode surface, prior to the o.r.c., was in a fully reduced state.All potentials in the study are quoted relative to a saturated calomel electrode (SCE). Variation in SERS intensity have been correlated in earlier studies with the extent of morphological and compositional damage by the focused laser., Under steady-state laser-damage conditions, as was observed for Ag/pyridine SERS, reproducible spectra were obtained. Intensity variations were therefore monitored directly in the present study.M. R. MAHONEY AND R. P. COONEY 2125 I I I I I I 1 I I I J 2000 2100 2200 Raman shiftlcm-' Fig. 2. Raman spectra from the surface of a polycrystalline copper electrode in aqueous 0.1 mol dm-3 Na,SO,, 0.1 mol dm-3 KCN. All spectra were recorded after an 0.r.c. : (a) -0.6, (b) -0.8, (c) -1.0, ( d ) -1.2 and ( e ) -1.4V.Table 1. Proposed assignments of v(CN) sub-bands" Cu/CN- solutions : previous assignments SERS [CN-]/[Cu'] 5 4 /crn-l /ern-' cm-I species ref. 2077 2077 2077 CN- 1 1 209 1 2095 2094 [Cu(CN),'J3- 12, 13 2102 2108 2109.5, 2096 [Cu(CN),I2- 12 21 14 21 13 2 1 13 [Cu(CN),]- 12 2130 2137 2137 [Cu(CN),]- (?) 1 1 a Shoulders are partly resolved sub-bands which vary intensity with changing potential or concentrations. RESULTS AND DISCUSSION POTENTIAL-STEP 0.R.C . Previous investigations of the copper/cyanide electrochemical SERS system have employed a multistep ramp-hold 0.r.c. in order to observe the intense Raman spectr~m.~ A less complex potential-step 0.r.c. was developed in the present study as a result of two series of experiments.First, the electrode potential was held for constant time intervals at a series of anodic limits. Secondly, the electrode potential was held at a series of anodic limits until an approximately constant amount of oxidizing charge (ca. 40 mC cm-2) was passed. In both series of studies SERS was2126 SERS BY CYANIDE ON COPPER ELECTRODES \ \ \ \ \ 20 40 60 80 100 tlmin 0 1 start 0.r.c. Fig. 3. Effect of laser interruption in the intensity against time curve for the line at ca. 2105 cm-l from a copper electrode surface. The potential before and after the 0.r.c. was - 1.25 V. For the dotted part of the curve the electrode surface was not under laser illumination. observed only after an 0.r.c. to +0.4 V. Overall, the preferred potential-step 0.r.c. is as follows: - 1.4, - 1.25 or -0.6 V e + 0 .4 V with an anodic pause of 5 s. APPEARANCE OF THE SERS SPECTRUM Two potentials (- 1.25 and -0.6 V) were selected for the majority of studies of SERS by cyanide on copper electrodes. The SERS spectrum at - 1.25 V after an 0.r.c. includes features at 21 13 cm-l [v(CN)], 1645 cm-l (H,O, v,) and 988 cm-l (Sot-, vl). The latter two features arise Trom the bulk electrolyte and exhibit comparable intensities before and after the 0.r.c. The SERS line (ca. 21 10 crn-l, with precise position depending on potential) is broad, structured and asymmetric, especially at -0.6 V (see fig. 1). Identification of the potential of maximum SERS intensity is complicated by a time-dependent decay (fig. 1). Nevertheless, it is clear that the intensity of the v(CN) band does vary with potential (fig.2). At potentials more cathodic than - 1.2 V the SERS spectrum becomes weaker and at very cathodic potentials (e.g. - 1.6 V), a loss in v(CN) band intensity occurs and can only be reversed by subjecting the working electrode to a further 0.r.c. Furthermore, the v(CN) band profile becomes sharper and shifts to lower wavenumbers as the potential is adjusted in a cathodic direction (fig. 2). This change in band profile appears to result from changing relative intensities of component sub-bands evident in the spectrum at -0.6 V (see fig. 1) and at other potentials. Some spectra appear to incorporate four or even five shoulders, which are assumed to arise from sub-bands within the v(CN) band profile. The positions and probable assignments of these sub-bands to cyanocopper(1) complex species (see below) are given in table 1 .INDICATIONS OF LASER DAMAGE After a potential-step o.r.c., the SERS v(CN) line was observed to decay with time. This decay has also been reported by Benner et aZ.7 In order to test for laser-induced dissolution processes, laser-interruption studies of the type previously reported3 for Ag/CN- were performed. Interrupting laser illumination of the working electrode surface shortly after the SERS intensity had attained a maximum at -0.6 V (i.e.M. R. MAHONEY AND R. P. COONEY 2127 ca. 10 min after the 0.r.c.) did not appear to modify the rate of intensity decay substantially. However a similar experiment at - 1.25 V did reveal an upward displacement of the intensity against decay curve (fig.3) following interruption of laser illumination of the surface. Therefore it appears that at - 1.25 V the laser was assisting dissolution of the intensely-scattering phase. In the light of the uniform-flux data (below) the absence of a laser-interruption ' step' at - 0.6 V probably reflects a steady state between laser-assisted oxidation of the surface and laser-assisted dissolution effects at anodic potentials. The difficulties of relating the entire surface sampling of voltammetry with the laser microzone (ca. 0.04mm diameter) sampling of SERS has been overcome in other systems3 by investigating uniform-laser-flux perturbations of the cyclic voltammogram. In these experiments, unfocused laser flux was employed to fill the working electrode surface (the 1 mm diameter cross-section of a copper wire in the uniform-flux studies) while cyclic voltammograms were recorded.For a cyanide-free 0.1 mol dm-3 Na,SO, electrolyte the introduction of uniform laser flux resulted in a slight increase in peak current at +0.4 V, which is presumed to be associated with increased generation of copper(rr) specie^.^ Uniform-flux perturbations for the Cu/O. 1 mol dmP3 Na,SO,, 0.01 mol dm-3 KCN system included the appearance of a shoulder on the reduction peak at - 0.05 V [tentatively attributed to unstable cyanocopper(rr) complexes]. Generally, all of the waves sensitive to cyanide concentration revealed an increase in current amplitude resulting from incident uniform laser flux. The observation of voltammetric perturbations at ca.0.1% of SERS flux (i.e. an unfocused laser) suggests3. that at SERS (focused) laser-flux levels such perturbations would be substantial. The indications of laser damage are similar to those previously reported for the Ag/CN- SERS system.,-, A similar explanation3 would be expected to apply to the Cu/CN- SERS case, i.e. the formation of a localized warm surface-zone (possibly as a result of photon-surface resonance) and the generation of convection currents in the interfacial region would be expected to accelerate both metal-surface corrosion and corrosion-complex dissolution. POTENTIAL- AND pH-DEPENDENCE STUDIES Although the 0.r.c. is expected to involve the generation of copper(rr) at the anodic limit (+ 0.4 V), the structured v(CN) band profile of the SERS spectrum incorporates sub-band components identifiable with cyanocopper(r) complexes (table 1).The lower the v(CN) band the higher the coordination number of the cyanocopper(r) complex. Therefore, the decrease in v(CN) as the potential is adjusted in a cathodic direction (fig. 2) would arise from an increase in the ratio [CN-]/[CuI] associated with the reduction of Cu* to copper metal at those potentials. Also, the highest v(CN) observed, at -0.6 V, coincides approximately with the upper limit (2137 cm-l)ll observed for cyanocopper(r) complexes. The pH dependence of cyanide SERS from copper electrodes is significantly different from that for silver electrodes in the same e1ectrolyte.l. l4 The extinction of Ag/CN- SERS at pH ca. 7 was interpreted (together with other evidence) as indicating the involvement of silver oxide in the intensely-scattering phase.'* 2, 4 9 l4 The SERS spectrum from copper/cyanide could be observed for a stirred electrolyte at pH values as low as 2.5.Starting with an electrolyte pH of 10.5, the SERS spectrum was generated using an 0.r.c. and then sufficient 1 mol dm-3 H,SO, was added to lower the pH to 3.5 and the solution stirred for 10 min. At the end of this time the Raman spectrum revealed no v(CN) bands. However, after a second 0.r.c. in the electrolyte at pH 3.5 the 2095 cm-l SERS band was once again observed, with intensity2128 SERS BY CYANIDE ON COPPER ELECTRODES J I I I I I I I I I 1 2000 2100 2200 Raman shiftlcrn-’ Fig. 4. Effect of electrolyte pH on the Raman spectrum from a copper electrode in aqueous 0.1 mol dm-3 Na,SO,, 0.01 mol dm-3 KCN.Electrode potential: -0.6 V. Spectra recorded after an 0.r.c. (a) pH 10.5, (b) pH 3.8 and (c) pH 2.3. Both (b) and (c) were recorded at 1.5 x the sensitivity of (a). comparable to that recorded at pH 10.5. A similar result was obtained after an 0.r.c. at pH 2.5. However, if the pH was lowered to a value < 2 then after an 0.r.c. no SERS was detected. Similar pH-dependence results were obtained for a parallel study using an unstirred electrolyte. In that case, lowering the pH from ca. 10.7 to 3.5 resulted in a minor change in the intensity of the v(CN) band. The copper electrode was then reduced at - 1.6 V and on returning to - 1.25 V no v(CN) band was detected. After a further 0.r.c.at the lower pH (3.5) the v(CN) band reappeared with comparable intensity to that observed prior to the reduction of the surface. Similar results were obtained when the electrolyte pH was adjusted to ca. 2.2. The pH dependence of SERS intensity at -0.6 V was also investigated using a different method. The pH of the electrolyte was adjusted to the value of interest prior to inserting the copper electrode. A freshly prepared copper electrode was used at each electrolyte pH value (fig. 4). In this case the spectra at pH ca. 2.3 were considerably weaker (see fig. 4). Examination of the potential against pH diagram for the copper/water system15 shows that at pH < 2 copper(1) oxide would not be expected to form. Copper(I1) oxide is unstable at pH < 4. As with the silver/cyanide system, the pH of extinction of SERS coincides approximately with the pH at which all oxide phases become unstable.However, the pH stability of copper(1) cyanide complexes complicates an interpretation based on oxide involvement. The pH stability of such complexes was studied using a series of solutions of CuCN in aqueous KCN having [CN-]/[Cu*] ratios 5 4. The Raman spectra of these solutions in the v(CN) region suggested the presence of an equilibrium of complexes (table 1). The pH stability of these complexes was investigated by observing changes in the partly resolved v(CN) profile of the Raman spectra as the pH was adjusted by the addition of dilute aqueous H,SO,. The spectrumM. R. MAHONEY AND R. P. COONEY 2129 of the [CN-]/[Cul] z 4 solution at pH cu.11 exhibited two intense v(CN) lines (2095 cm-l assignedl27 l3 to [Cu(CN),I3- and 2078 cm-l assignedll to CN-}. When the pH was lowered to ca. 4.6, CuCN [identified by its infrared v(CN) band at 2164 cm-lI6 was precipitated. The Raman spectrum of the supernatant solution exhibited intense lines at ca. 2098 cm-l ([CU(CN),]~-)~~~ l3 and ca. 21 12 cm-l ([Cu(CN),]-}12. Lowering the pH to 2 resulted in more precipitation and the appearance of a single band (2096 cm-l, [CU(CN),]~-}~~. l3 in the spectrum of the solution phase. The intensity of the v(CN) band profile at pH 2 is substantially less than the intensity at pH ca. 11. Therefore the observation that copper/cyanide SERS is undiminished at pH ca. 3.8, i.e. below the onset (pH > 4.6) of substantial cyanocopper(1) complex decomposition, does suggest that oxide stabilization is involved in the intensely-scattering phase. The point of SERS extinction (pH ca.2) also coincides with the lower limit of copper oxide stability. It is possible that the apparent coincidence of v(CN) components in the SERS spectrum (table 1) with lines attributed to cyanocopper(1) complexes is accidental and that the intensely-scattering phase is an oxidesyanide phase (as was found for Ag/CN- SERS).' However the apparent appearance and disappearance of v(CN) SERS components with changing potential (above) is indicative of a cyano complex equilibrium. CONCLUSIONS SERS by cyanide on copper electrodes appears to arise from an equilibrium mixture of cyanocopper(1) complexes, probably stabilized by association with copper(1) oxide within the laser microzone.This compares with Ag/CN- and Au/CN- SERS, which appear to have an oxide-cyanide compound as the intensely scattering phase.'? l6 It is possible that the involvement of oxide-cyanide components in these SERS phases arises because the oxide associated phase is more adherent and less soluble than simple cyanometal complexes under the conditions of photochemical warming within the laser-damage microzone. Also, if laser-assisted corrosion of the metal surface is important (as has been found for Ag/CN- SERS and is suggested by data presented here), then laser-assisted diffusion of soluble cyanometal complexes out of the laser microzone may deplete the microzone of cyanide. At the high pH of the electrolyte such cyanide depletion may favour the formation of oxide-cyanide species within the laser-illuminated zone of the surface.Localized oxygen-rich laser-damage zones (ca. 0.13 mm diameter) have been previously identified4 for Ag/CN- SERS. As a result of these conclusions, Cu/CN- SERS is classified as type I11 (inorganic microzone).2 We thank the Australian Research Grants Scheme for supporting this project. M. R. Mahoney and R. P. Cooney, Chem. Phys. Lett., accepted for publication. R. P. Cooney and T. P. Mernagh, J. Electroanal. Chem., 1984, 168, 67. M. R. Mahoney and R. P. Cooney, J. Phys. Chem., 1983,87,4589. R. P. Cooney, T. P. Mernagh, M. R. Mahoney and J. A. Spink, J . Phys. Chem., 1983, 87, 5314. T. P. Mernagh, R. P. Cooney and J. A. Spink, J. Raman Spectrosc., 1985, 16, 57. G. Laufer, T. F. Schaaf and J. T. Huneke, J . Chem. Phys., 1980, 73, 2973. ' R. E. Benner, K. U. von Raben, R. Dornhaus and R. K. Chang, Surf. Sci., 1981,102, 7. B. H. Loo, Spectrosc. Lett., 1982, 15, 85. M. R. Mahoney, M. W. Howard and R. P. Cooney, Chem. Phys. Lett., 1980,59, 71. C. Kappenstein, R. P. Hugel, A. J. P. Alix and J. L. Beaudoin, J. Chim. Phys., Phys-Chim. Biol., 1978, 75, 427. lo D. A. Long, Raman Spectroscopy (McGraw-Hill, New York, 1977), p. 134. l2 M. J. Reisfeld and L. H. Jones, J . Mol. Spectrosc., 1965, 18, 222. 70 FAR 812130 SERS BY CYANIDE ON COPPER ELECTRODES l3 G. W. Chantry and R. A. Plane, J. Chem. Phys., 1961, 35, 1027. l4 M. R. Mahoney and R. P. Cooney, J. Raman Spectrosc., 1981, 11, 141. l5 U. Bertocci and D. R. Turner, in Encyclopedia of Electrochemistry of the Elements, ed. A. J. Bard l6 M. R. Mahoney and R. P. Cooney, J. Chem. SOC., Faraday Trans. I, 1985, 81, 4/1885. (Marcell-Dekker, New York, 1974), vol. 11. chap. 11-6. (PAPER 4/ 1886)

ISSN:0300-9599    

DOI:10.1039/F19858102123

出版商:RSC    

年代:1985

数据来源: RSC

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J. Chem. SOC., Faraday Trans. I , 1985,81, 2131-2144 Infrared Spectroscopic Studies of the Solvation of Aprotic Solvents and Ions in Methanol? BY HOWARD L. ROBINSON AND MARTYN C. R. SYMONS* Department of Chemistry, The University, Leicester LE 1 7RH Received 15th November, 1984 The effect of the addition of various polar aprotic solvents (B) and electrolytes on the first overtone of the 0-H stretching band (2v,,) of CH,OH in CD,OD or CD,OH in CD,OD has been measured. The results for the solvent mixtures are discussed in terms of contributions from MeO-H * * * B oscillators and terminal methanol molecules, designated lone-pair free [(LP),,,,] units. For the electrolytes, in addition to contributions from methanol molecules bound to the cations and anions, we consider contributions from (LP)free groups and from other weakly bound methanol molecules, including possible contributions from non-hydrogen-bonded OH groups, (OH)free. The results are compared with recent proton magnetic resonance studies and with similar studies of aqueous solutions.It is concluded that basic aprotic solvents and anions form fewer hydrogen bonds to methanol molecules than to water molecules. For example, dimethyl sulphoxide forms one hydrogen bond in methanol (two in water), triethyl- phosphine oxide forms two hydrogen bonds in methanol (three in water) and chloride ions form four hydrogen bands in methanol (six in water). For some time changes in the optical spectra of solvent and solute molecules have been used to probe the structures of solutions, especially those involving protic solvents [see, for example, ref.(1)-(91. More recently, the techniques of diffraction and computer simulation have been developed and they provide results with which our own conclusions can be compared. In much of our work we have selected methanolic rather than aqueous systems because many problems associated with water structure are absent and yet the modes of interaction with solutes are generally comparable. In our view,6 when basic aprotic solvents (B) or anions (A-) are added to methanol, the following equilibria dominate :6 Me Me Me Me For ease of representation, B is depicted as having a solvation number of one in eqn (1). It may, of course, have a solvation number greater than one, as is usually the case for anions. We therefore expect to find spectroscopic evidence for two new units.The unit Me Me I 1 B.. .HO.. .HO.. . t Taken as Solvation Spectra, Part 75. 2131 70-22132 I.R. STUDY OF SOLVENTS AND IONS IN CH,OH represents the aprotic solvent, B, hydrogen bonded into a chain of methanol molecules. We expect the 0-H oscillator of the molecule directly bound to B to be significantly modified, the second and subsequent molecules being, to a first approxi- mation, indistinguishable from normal, ‘ bulk ’ methanol, (MeOH),,,,. Similarly, the terminal MeOH molecule in the (LP)free unit should be distinguishable but the other molecules of this unit are expected to resemble bulk methanol. We have previously suggested that absorption bands at ca. 3440cm-l in the fundamental (voH) and ca.6790cm-l in the overtone (2v0,) are due to (LP)f,,e terminal molecules.6’ Monomeric methanol molecules have a relatively narrow 2v0, band at ca. 71 15 cm-l and (OH)free units, i.e. terminal non-bonded (OH) groups Me Me I I HO ...HO... are also expected to absorb in this region. However, when O-H-.-X units are formed, the band broadens and shifts to low energies, the average ‘bulk’ methanol peak being at ca. 6350 cm-l in the first overtone. The hydrogen bonds for average di-bonded methanol molecules are stronger than those formed by the terminal molecules in (LP)f,ee units because of the reinforcing effect of the bond to oxygen. Since the shift to low frequency is approximately proportional to the hydrogen-bond strength, we conclude that the hydrogen bonds formed by terminal molecules are slightly greater than half the strength of the average bond.When cations (M+) are added, we can depict the solvation process by a series of equations of which the first is Me Me Me Me We define the ‘solvation number’ of an ion or aprotic molecule as the number of solvent molecules interacting directly therewith. Hence, in this case, the number of M+-..O bonds formed by the cation is its solvation number. Note that this is by no means the only definition of ‘ solvation number’.8 units to remain as such, since methanol molecules have two lone pairs of electrons able to accept hydrogen bonds. Thus the reaction Me Me Me Me Me Me In this case, however, we do not expect the I l l I l l (3) . . .OH.. .OH.. .OH.. . + (OH)free + . . .OH...OH.. .OH.. . H Me-0- - - is expected to scavenge most of the units. Of the three hydrogen bonds formed by the cross-linking methanol molecule, those to oxygen will be weaker than the average bonds, whilst that involving the 0-H proton is expected to be stronger than average. Possible contributions from these units to the infrared spectra are considered in the Discussion. In a previous infrared study of methanolic solutionsg we established that, at ambient temperatures, the fundamental band (vOH) was of little use for extracting structural information. However, on cooling to between - 70 and - 100 “C the bands narrowed and shifted and, in several cases, resolution into two or more components was possible.H. L. ROBINSON AND M. C. R. SYMONS 2133 Because of experimental difficulties no attempt was made to quantify these results.However, we have recently shown that there is a linear correlation between infrared shifts in vOH and proton resonance shifts (OH)1° and this was used to estimate a range of separate ion shifts and solvation numbers, due allowance being made for shifts caused by (LP)f,ee units. Our present aim was to attempt to obtain quantitative information about cosolvent- and salt-induced changes in the 2v0, region in the hope of confirming and extending the n.m.r. studies.1° Furthermore, the marked contrast in spectral behaviour of methanol and water in this region is thought to be of some significance. EXPERIMENTAL All the first overtone methanol spectra (1650-1300 nm) were recorded on a Perkin-Elmer 340 u.v./vis./n.i.r.spectrophotometer. Spectra were recorded at 25kO.l "C and 50f0.1 "C on a photometric scale of 0-0.2, a 1 expansion of five-fold and a 3 min scan time. All salts used were of AnalaR grade and were dried under vacuum and stored in a dessicator over silica gel before use. Cosolvents, of the best available grades, were purified by distillation over CaH, and dried over CaH,. Methan[2H,]ol, 99.5 % , was used without further purification. Each sample contained 20% MeOH by volume. A known volume of cosolvent or salt + MeOD solution was added and finally MeOD was added to make up to the 5 cm3 mark. The reference contained the same volume of cosolvent or salt+MeOD solution and MeOD was added to make up to the 5 cm3 mark. Stock solutions of MeOD+salt were prepared by weight so that the mole fraction of salt was established.It is useful to keep the concentration of OH oscillators constant in the sample solutions so that spectral changes are only a result of salt or base effects on the solvent and not caused by an imbalance of OH groups throughout the systems studied. Glass cells with a 2 mm pathlength were used. Absorbance due to the cells was recorded with the baseline correction before each sy s tem studied. RESOLUTION-ENHANCEMENT PROCEDURE A resolution-enhancement technique was applied to methanol (CH30H or CD30H) in CH,OD (or CD,OD) at 25 "C. Data coordinates were read directly into a 280 microcomputer from the spectra. The resolution enhanced spectra was obtained by subtracting a fraction ( k ) of the second-derivative spectrum from the original absorbance, the fraction (k) of the second derivative being changed until a good resolution enhanced spectrum was obtained [see fig.1 (b)]. This procedure, developed in these laboratories, is qualitatively useful for establishing the presence of poorly resolved bands. The resulting spectra have not been used in our curve analyses. CURVE-ANALYSIS PROCEDURE The first overtone methanol/MeOD spectra were deconvoluted using computer graphics. The experimentally obtained spectra, which were recorded in nanometers on a Perkin-Elmer 340 spectrophotometer, were first converted to wavenumbers in order to remove any skew effect that the nanometer scale might have on the bands. In all cases we used the minimum number of symmetrical bands required to reproduce the experimental spectra within experimental error.All bands gave rise to shoulders in the spectra and hence were not arbitrarily selected. The spectral envelope could be simulated with a minimum of four bands. This was achieved using a standard VIDCA program (visually displayed curve analysis). This program is a series of subroutines and functions, written in FORTRAN, which can produce and alter one or more curves with variable heights, positions, half-widths and Gaussian-Lorentzian properties, displaying them separately, as well as combined, on a visual display unit. The core program uses the following algorithm to compute the approximate band shape:2134 0.: a2 c * 0.1 0 4 9 I.R. STUDY OF SOLVENTS AND IONS IN CH,OH 1494 nm bulk methanol (1575 nm) k hydrogen bonds (1435 nm) entities (1410 nm) I I I I I I I )o 1350 1400 1450 1500 1550 1600 1650 wavelength/nm wavelength/nm Fig: 1.(a), (b). For description see opposite. where X, is the band height in absorbance units, X , is the band position in abscissa units, X, = (2l-"- l)1'2/n, X, = [a ln(2)]:/n, v is the abscissa value at which absorbance is to be calculated, n is the half-bandwidth at half-height, a is a Gaussian contribution, B is the baseline in absorbance units and rn is the number of bands producing the overall envelope. The computed spectrum envelope is obtained by combining each band found by the above equation : rn i-1 Ycalc = & where Ycalc is the computed absorbance fit and & is the calculated ordinate for each band. Using this procedure, the individual band parameters were altered until a good fit with the experimental band was obtained.Copies of the final deconvoluted spectra were obtained using a CALCOMP 1012 plotter.H. L. ROBINSON 0 175 al c 2 0.1 2 .o 0 025 AND M. C. R. - - SYMONS 21 35 'i I wavenumber/cm-' 2t // F / O / L,o'O , I , 0 20 40 60 T I T Fig. 1. Infrared spectrum of methanol (10% CH,OH in CH,OD) in the 2vOH region. (a) As a function of temperature [(l) 25, (2) 30, (3) 35, (4) 40, (5) 45, (6) 50 and (7) 55 "C]. (b) At 25 "C, using a resolution-enhancement procedure [this is no longer a true spectrum, but serves to help identify the presence of the main bands selected in (c)]. (c) As fitted using bands I-IV (the high-frequency tail has been omitted since it is not due to 0-H oscillators in the first overtone region, otherwise the sum of the individual peaks effectively reproduces the experimental spectrum).( d ) Growth of bands 1-111 as a function of temperature (error bar shown) (A, I; 0, 11; 0,111)- RESULTS AND DISCUSSION PURE SOLVENT Spectra in the 2v,, region for a 10% solution of CH,OH in CH,OD at a range of temperatures are shown in fig. l ( a ) . The most striking aspect is the long high-frequency 'tail' in the region of weakly bound 0-H groups, with a small step at the onset of absorption, which is the value of ca. 7100 cm-l (1408 nm). [That there are intermediate bands between the 'free' and strongly bound (6350 cm-l) (1 575 nm) regions is more clearly demonstrated by the resolution-enhanced spectrum, shown in fig.1 (b).] The experimental spectrum can be reconstituted with a minimum of 4 symmetrical bands, as indicated in fig. 1(c). There is no point in using more, though this could, of course, be done. We stress that this fit does not prove the presence of 4 bands due to 4 distinct types of oscillator. In particular, the ' onset' band (band I) at 7100 cm-l may be part of a skew band with a maximum at ca. 6935 cm-l (band 11) and there may also be weak contributions from other combination bands.ll However, the spectrum due to CD,OH in CD,OD was very similar in this region, so we do not consider that such bands are of major importance. In the following we generally treat these bands as one feature (I + 11) which we assign to 'weakly bound' 0-H oscillators (OH),,, which is presumed to include the weakly bound oscillators present in the branched-chain units shown in eqn (3), thought to be formed from genuine (OH)free units.21 36 I.R.STUDY OF SOLVENTS AND IONS IN CH,OH Table 1. Data obtained from the (LP)free band for methanol gradient for deconvoluted cosolvent gradient ( LPhr ee band solvation no.a Et,PO HMPA DMSO THF DMA CD,CN BuiNCI Et,NCI Me,NCl BuiNBr Et,NBr BuiNI BuiNCH,CO, Et,NCN Et,NOCN BuiNNO, BuiNCIO, DMSO-d, 0.63 0.48 0.38 0.36 0.35 0.34 1.16 1.17 1.10 1.12 1.04 1.04 1.01 0.62 0.63 0.61 - 0.094 - - 0.05 1 - - 0.045 0.186 - - - 0.082 2 1 1 1 1 < 1 4 4 4 4 6 4 6 4 < 4 3 2 3 2 3 2 6 2 2/ 1 a Estimated from the gradients assuming a solvation number of 2 for Et,PO. Trends in the intensities of bands I, I1 and I11 on changing the temperature are shown in fig.1 (d). As expected, on heating there is a general increase in the intensities of the bands for more weakly bound oscillators at the expense of the strongly bound oscillators (band IV). Band 111, with v,,, at ca. 6715 cm-l, is close to a band which develops strongly on the addition of basic cosolvents, which we previously identified as due to (LP)f,ee or terminal methanol molecules.6 Clearly there must be contributions from other types of hydrogen-bonded methanol in this region but we suggest that changes in this band on adding solutes largely reflect changes in the concentration of (LP)free units. Further justification for this assignment stems from the results described below and we use this description throughout.Less weight is given to changes in the main band (band IV) at ca. 6350 cm-l because of contributions from intermolecular coupled bands of the type studied by Berneau and Corset12 and from the high-frequency edge of other bands. Only for the weak bases, cyanomethane (CH,CN) and Cloy, was it necessary to introduce an extra feature in these reconstructions. Detailed results for changes in the ‘deconvoluted’ (OH),, and (LP)free bands were obtained for a selected number of solutes (tables 1 and 2). These analyses yielded similar relative trends to those obtained by measuring changes in absorbance at the band maxima. ADDED APROTIC SOLVENTS Here we focus attention on eqn (1). Since there are few, if any, genuine groups, we need not, to a first approximation, considerH.L. ROBINSON AND M. C. R. SYMONS 2137 Table 2. Infrared parameters for deconvoluted band maxima in the 2v0, region for methanol at 25 “C position band nm cm-l (1) weak hydrogen-bonds 1408 7100 (2) weak hydrogen-bonds 1442 6935 (3) ‘ ( W f r e e ’ band 1487 6725 (4) bulk hydrogen-bonds 1579 (1554)u 6335 ClO;*..HOMe 1429 7000 CD,CN...HOMe 1449 6900 Value for CD,OH in CD,OD. which in our view dominates the behaviour of water under these condition^.^'^^ However, we expect, and find, a marked growth in the (LP)f,ee band at ca. 6725 cm-l. This is shown for the ‘deconvoluted’ band in fig. 2 for MeOH+DMSO and MeOH + Et,PO. The major features to note are: (i) for these relatively strong bases, there was no need to postulate the growth of an extra band in the high frequency region and hence the B..-HOMe units are presumed to contribute to the broad ‘bulk methanol’ band, (ii) there is always a marked growth in the (LP)free band, (iii) its rate of growth varies with the nature of the added solvent and (iv) there is a small loss of the (OH),, band( s) .The slopes of the graphs shown in fig. 2 and related plots give the growth rate of (LP)f,,e groups as solute is added. These data are then converted into approximate solvation numbers of the solute, using eqn (1) together with the following calibration. It isconvenient to use as a reference the results of our detailed study of triethylphosphine oxide (Et,PO) in methanolic systems,14 which established that in methanol-rich solutions each molecule forms two hydrogen bonds. If we link the initial slope of the plot for Et,PO with a solvation number of two, the solvation numbers for the other solvents result from direct comparison.These are given in table 1. For dimethyl sulphoxide (DMSO), dimethylacetamide (DMA) and tetrahydrofuran (THF) the slope is ca. half that for Et3P0, implying that the predominant solvation number for these bases is one in dilute methanolic solutions. This result accords well with our estimate for DMSO and DMA using the S-0 and C=O stretching bands.15 For HMPA the results suggest that both mono- and di-solvates are formed. However, our infrared studies using the P-0 stretch suggest that the latter d0~ninates.l~ We have no corroborating evidence for THF but would not have expected this relatively weak base to form more than a mono-solvate in methanol.There are small decreases in the (OH),, region, as expected, and also in the main band at ca. 6350 cm-l. The latter decreased more rapidly when deuterated solvents were used, indicating contributions from combination bands for the normal (lH) solvents,12 so no attempt has been made to quantify the changes. For cyanomethane (CH,CN or CD,CN) a different approach was needed because the solvate Me MeCN...HO... 12138 I.R. STUDY OF SOLVENTS AND IONS IN CH,OH i concentration of solute (mole fraction) Fig. 2. Changes in band I11 for methanol on the addition of a range of solutes (error bar shown): A, BuiNfI-; A, Et,N+Cl-; 0, BuiNTlO;; 0, Et,PO; 0, [2H,]DMSO; +, CD,CN. has a relatively weak bond so the HO oscillator contributes in the high-frequency region of interest.To obtain a good fit for the spectra an extra feature at ca. 6900 cm-l had to be included. The overall changes in the five components are shown in fig. 3. Although the accuracy of our analysis must be reduced by the need to include this band, the changes in the normal methanol bands are as expected. In particular, there is a gain in the (LP)free region suggesting a solvation number d 1. We conclude that a solvation number between 0 and I is required for MeCN in methanol, and this is in good accord with our studies of the v(CN) band. We conclude that most basic aprotic solvents prefer to form one hydrogen bond to methanol in dilute solutions. The strong bases Et,PO and HMPA probably form two bonds, whilst the weak base MeCN is, in part, not hydrogen bonded.ELECTROLYTE SOLUTIONS We have adopted the same procedures for electrolyte solutions. In previous worklo we have established that proton resonance shifts induced by tetra-alkylammonium ions, R,N+ (except where R is CH,), are close to zero. Here we argue that these ions are unlikely to form hydrogen bonds to the available oxygen lone pairs of electrons and hence their only influence is likely to be one of dilution. In previous proton resonance studies6* lo we have shown that this dilution effect does not greatly modify the total hydrogen bonding in methanol. As a further check on this postulate we haveH. L. ROBINSON AND M. C . R. SYMONS 2139 L o'2 t (4 /// 4 \\\ \ w 1 4 9 2 nm loss of bulk MeOH wavenum ber/cm-* wavelength/nm o.zo/ O \ t 0 16 t '0 -0-• I 0 0-0-0 0 4- 0-0 I I I I O O 0 04 0.12 concentration of CD,CN (mole fraction) Fig.3. Changes in the spectrum of methanol on the addition of CD,CN. (a) On increasing the mole fraction of CD,CN [( 1) 0.00, (2) 0.086, (3) 0.129 and (4) 0.1721, (6) showing the need to include band V in the synthesis of curve (4) and (c) trends in these bands (0, I; 0, 11; 0, 111; 0, IV; A, V). studied -the effect on the near-infrared spectrum of MeOH in MeOD of the addition of tetramethyl silane. Spectral changes are small compared with those induced by R,N+X- salts, so we conclude that these can indeed be assigned to anion solvation only. This is further supported by the fact that the general form of the spectral changes in the (OH),,, (LP)free and bulk MeOH regions are very similar to those induced by basic aprotic solvents, except for R,N+ClOy salts, whose behaviour resembles that2140 I.R.STUDY OF SOLVENTS AND IONS IN CH,OH of MeCN (see below). The key results are a rapid increase in band I11 assigned to (LP)free groups (fig. 2) with a fall in the ‘bulk’ band and a small decrease in the (OH),, features. Again, to a first approximation, we consider it reasonable to assign these large changes in the (LP)free band to the gain of these groups according to eqn (l), where B is now the anion. The number of (LP)free groups generated per anion should approximately equal the solvation number of the anion. By comparing the slopes of the linear regions of the plots shown in fig. 2 with that for our reference (Et,PO), we obtain solvation numbers of ca.4 for C1-, Br- and I-. Other results are given in table 1. In our previous, completely independent, analysis of n.m.r. shifts induced by halide ionslO we also obtained solvation numbers of ca. 4 for these ions. This good agreement strongly supports both procedures and we have considerable confidence in the result. The strongly basic acetate ion is also found to have a solvation number of 4, presumably each oxygen forming two hydrogen bonds (structure I). However, the linear ions CN- and NCO- have solvation numbers of 2. We suggest that these bonds are also linear, as in structure 11. , HO-Me 6 , HOMe / - The more weakly basic nitrate ion also has a solvation number of 2, although one might have expected 3, with one bond to each oxygen.However, the nitrate ion has a high polarisability and we suggest that solvation at two oxygens pulls the negative charge away from the third oxygen. This result accords well with our conclusion, based on an infrared study of nitrate ions, that these ions are asymmetric in dilute methanolic solution in the absence of significant ion pairing.17 As is well known, perchlorate salts give rise to the gain of a resolved band in the (OH),, region.1s-20 There has been some controversy regarding the origin of this feature, some favouring the idea that it is due to (OH)free units produced by a ‘structure breaking’ influence and others that they are due to units, the hydrogen bonds to C10, ions being very weak. We favour the latter de~cription.~3 1 9 9 2o Using this concept, we have introduced a fifth band at ca.7000 cm-l assigned to oscillators (fig. 4). Changes in the other bands agree with expectation. Thus band IV, from strongly bonded methanol falls, but band 111 grows, because of the gain in (LP)free groups. Bands I and 11, assigned to weakly bound methanol molecules, fall slightly, as observed for the basic aprotic solvents. Solutions of sodium tetraphenylboron behaved similarly (fig. 5). In this case, we assume that there is no appreciable hydrogen bonding to the BPh, ion (or that thisH. L. ROBINSON AND M. C . R. SYMONS 0 x 4 0.08- 0 2141 c concentration of But,NC10, (mole fraction) Fig. 4. Trends in the intensities of the deconvoluted ClO; bands for solutions of Bu,N+C10; in methanol (0, I ; 0, 11; 0, 111; 0, IV; A, V', an extra band assigned to ClO;--*HOMe solvates.c o----o-o--~ t L OO 0.02 concentration of NaBPh, (mole fraction) Fig. 5. Trends in the intensities of the deconvoluted bands for solutions of Na+BPh, in methanol (0, I ; 0, 11; 0, 111; 0, IV). bonding is extremely weak), and hence that the observed changes are largely induced by the sodium ions. Thus eqn (3) should dominate, with the generation of units which will react according to eqn (4) to give the (OH),, units. This is indeed observed. Note that there is also a loss of absorbance in the (LP)f,ee region. This supports the postulate that at 25 "C there is a significant concentration of (LP)f,ee groups, or terminal methanol groups, in liquid methanol at 25 "C.2142 I.R. STUDY OF SOLVENTS AND IONS IN CH,OH CATION SOLVATION As stressed above, we assume that the effects of R,N+ ions on the methanol spectra are small, as in the case for tetramethylsilane and as was found in our n.m.r.studies.lo Hence the effects for these salts are assigned entirely to the anions. However, for salts containing metal cations, eqn (2) should occur, giving rise to a gain in (OH),, units. This was observed experimentally for Na+BPh;. However, there was also a marked loss of absorption in the region of the (LP)free band, as is clearly illustrated after curve deconvolution (fig. 5). If we make the simple postulate that the initial, linear loss of (LP)free is a measure of the solvation number of sodium, and using as before the data for Et,PO as a reference, we obtain a solvation number of 5.7 for sodium in methanol.In our n.m.r. studies we obtained 5 + 1 for this value. Note that in a Monte Carlo statistical-mechanics simulation of Na+CH,O- in methanol, Jorgensen et a1.21 obtained a solvation number of six for Na+ ions. Even more interesting is that they ' see' a second layer of ca. 6 MeOH molecules. This is exactly what we imply when we write the unit as We have long maintained that the first solvent shell is bonded to other solvent molecules in this way3-5f10 and it is satisfying to have this idea supported by these computer simulations. Any attempt to use this procedure for the perchlorates must be treated with reserve, because we need to include an extra absorption band for ClO;...HOMe units. Nevertheless, by assigning a solvation number of 2 to C10; ions we obtain a value of ca.6 for Na+, in agreement with the results for NaBPh,. Similar treatment of the results for LiClO, gives a solvation number of ca. 4 for Li+, again in good agreement with the n.m.r. result (4).1° However, for Mg(ClO,), we estimate a solvation number between 7 and 8 for Mg2+, whereas this is known to be 6, at least for solutions at ca. - 70 0C.22 Our n.m.r. studies also gave a value of 6 for Mg2+ for solutions at 25 "C. In our view, this discrepancy is not significant and simply reflects the inaccuracies of the procedure used for cations. ION-PAIR FORMATION Our n.m.r. resultsl09 23 for tetra-alkylammonium ions together with earlier e.s.r. results for paramagnetic anions2* and low temperature resultsg strongly suggest that ion pairing should be extensive for concentrated R,N+X- solutions. This explains the marked curvature of the plots showing the growth of the (LP)free band (fig.2). We intend to treat these results quantitatively in the CONCLUSIONS Our major conclusions are that measurements of changes in the absorption spectra of methanolic solutions in the 2v,, region provide information leading to the solvation numbers of basic aprotic solvents and electrolytes. The resulting numbers accord well with those obtained from a previous analysis of n.m.r. data. Furthermore, results for dimethyl sulphoxide and cyanomethane accord well with those obtained from independent studies of the infrared spectra of these m o l e c ~ l e s . ~ ~ ~ l6 We conclude that these solvation numbers are probably reliable and should be of use in assessing the results of computer calculations on the structures of methanolic solutions.When these results are compared with those from aqueous solutions, it is clear thatH. L. ROBINSON AND M. C. R. SYMONS 2143 1300 1400 1500 wavelength/nm 1600 Fig. 6. Comparison of the near-infrared bands (2v,,) at 25 "C for (a) HOD in D,O and (b) CH,OH in CH,OD. water tends to form more hydrogen bonds to basic cosolvents and to anions, but solvation of cations is comparable in the two solvents. This is interpreted qualitatively in terms of the presence of a high concentration of groups in water but a negligibly low concentration of such groups in methanol. In contrast, both solvents have a relatively high concentration of (LP)free groups, which control the extent of cation solvation.Examples of the different solvation numbers for solutes in water and methanol are DMSO (water 2, methanol I), Cl- (water 6, methanol 4) and C10, (water 4, methanol 2). Finally, we draw attention to the remarkable difference between the 2v,, spectrum for liquid methanol (CH,CH in CH,OD) at 25 "C and that for water (HOD in D,O) at 25 "C (fig. 6). Symons and c o ~ o r k e r s ~ ~ ~ 24 and 26 have interpreted the narrow component occurring at the onset of absorption for HOD as being largely due to groups, but this view is by no means universally accepted. The alternative view is that liquid water contains no significant concentration of 'free' groups, but a large range of bonded groups, the strength of the bonds varying over a wide range.Two factors are thought to contribute to the initial peak. One is that the shift is initially not very sensitive to bond strength for very weak bonds and the other is that weakly bonded oscillators have greater oscillator strengths than strongly bonded oscillators in the overtone region. We agree with these arguments and have always suggested that there is a significant contribution from weakly bonded units. However, so far as we know, the same arguments apply to methanol. In this case, we agree that there are no significant concentrations of (OH)free units and that the high-frequency tail is indeed due to a range of weakly bonded groups. The marked contrast with the spectrum for water encourages us in the view that, for water, the concentration of (OH)free groups is high.We thank the S.E.R.C. for financial support, Dr Graham Eaton for continuous help throughout this study and Dr M. J. Blandamer for much helpful discussion.2144 I.R. STUDY OF SOLVENTS AND IONS IN CH,OH M. Smith and M. C. R. Symons, J. Chem. Phys., 1956,25,1074; Discuss. Faraday SOC., 1957,24,206. M . C. R. Symons, Philos. Trans. Roy. SOC. London, Ser. B, 1975, 272, 13. M. C. R. Symons, in Electron-Solvent and Anion-Solvent Interactions, ed. L. Kevan and B. Webster (Elsevier, Amsterdam, 1976). M. C. R. Symons, Pure Appl. Chem., 1981, 53, 223. M. C. R. Symons. Ace. Chem. Res., 1981, 14, 179. M. C. R. Symons, V. K. Thomas, N. J. Fletcher and N. G. M. Pay, J. Chem. SOC., Faraday Trans. 1, 1981,77, 1899. ’ M. C. R. Symons, N. J. Fletcher and V. K. Thompson, Chem. Phys. Lett., 1979,60, 323. B. E. Conway, Ionic Hydration in Chemistry and Biophysics (Elsevier, Amsterdam, 198 1). I. M. Strauss and M. C. R. Symons, J. Chem. SOC.. Faraday Trans. 1, 1977, 73, 1796; I. M. Strauss and M. C. R. Symons, J. Chem. Soc., Faraday Trans. I , 1978, 74, 2146. lo M. C. R. Symons, J. Chem. Soc., Faraday Trans. I , 1983, 79, 1273. l 1 C. Bourderon and C. Sandorfy, J. Chem. Phys., 1973,59, 2527. l2 A. Burneau and J. Corset, Can. J. Chem., 1973, 51, 2059. l 3 M. C. R. Symons, J. M. Harvey and S. E. Jackson. J. Chem. SOC., Faraday Trans. I , 1980, 76, 256. l4 G. Eaton and M. C. R. Symons, J. Chem. SOC., Faraday Trans. I , 1982,78, 3033. l5 G. Eaton, Ph.D. Thesis (Leicester University, 1983). l6 G. Eaton, A. S. Pena and M. C. R. Symons, unpublished results. T. J. V. Findlay and M. C. R. Symons, J. Chem. SOC., Faraday Trans. 2, 1976, 72, 820. G. E. Walrafen, J. Chem. Phys., 1970, 52, 4176. 67, 3435. l9 L. J. Bellamy, M. J. Blandamer, M. C. R. Symons and D. Waddington, Trans. Faraday SOC., 1971, *O M. C. R. Symons and D. Waddington, J. Chem. SOC., Faraday Trans. 2, 1975, 71, 22. *l W. Jorgensen, B. Bigot and J. Chandrasekhar, J. Am. Chem. SOC., 1982, 104, 4584. 22 J. H. Swinehart and H. Taube, J. Chem. Phys., 1962,37, 1579. 23 M. Krell and M. C. R. Symons, unpublished results. 24 J. Oakes, J. Slater and M. C. R. Symons, Trans. Faraday SOC., 1970,66, 546. 25 W. A. P. Luck, Discuss. Faraday SOC., 1967, 43, 1 15. 26 W. A. P. Luck, Angew, Chem., 1980, 19, 28. (PAPER 4/1939)

ISSN:0300-9599    

DOI:10.1039/F19858102131

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. 1, 1985, 81, 2145-2153 Kinetics of the Solvolysis of the trans-Dichloro tetrap yridinecobal t(m) Ion in Water + Dimethyl Sulphoxide Mixtures BY CHRISTINE N. ELCY AND CECIL F. WELLS* Department of Chemistry, University of Birmingham, Edgbaston, P.O. Box 363, Birmingham B15 2TT Received 20th November, 1984 The kinetics of the solvolysis of l,B-[Copy,Cl,]+ have been followed over a range of temperatures in mixtures of water with DMSO up to a mole fraction of DMSO(x,) z 0.2. The logarithm of the first-order rate constant varies approximately linearly with the reciprocal of the dielectric constant at the same temperature, contrary to findings for the solvolysis of this complex in mixtures of water with alcohols. Extrema in the enthalpy, AH*, and entropy, AS*, of activation for solvolysis of the complex in the latter mixtures occur at low x, of the cosolvent and correlate well with the extrema found in the relative partial molar volume, % - e, of the cosolvent.With additions of DMSO to water, sharp changes in 6- do not occur until x, z 0.2, and extrema for AH* and AS* also do not occur until x, z 0.2 is reached. The application of a free-energy cycle by combining free energies of activation, AG*, with free energies of transfer, AG:, shows that the effect of solvent structure is greater on the pentacoordinated cation in the transition state than on the hexacoordinated cation in the initial state. The kinetics of the solvolysis of the trmwdichlorotetrapyridinecobalt(rr1) ion have been investigated in water + sucrose mixtures1 and in a series of water +cosolvent mixtures2 where the cosolvent is methanol, propan-2-01 or t-butyl alcohol.For such a dissociative-type reaction, with extension of the Co..-Cl bond in an unstructured continuous medium, a plot of log k against D;' (where k is the first-order rate constant and D, the dielectric constant of the medium) should be l i ~ ~ e a r , ~ - ~ but in all the above cases2 curves are obtained. The expressions derived by Laidler and Landskroener3 on the basis of a continuous medium and the application of transition-state theory to the bond extension MZ~-.-XZx in the complex Czc can be modified to allow for the effect of solvent structure : + AG,"(MZ~), + AG:(XZx), - AG:(CZc), (1) where N is Avogadro's number, e is the electronic charge, r is the radius of the complex, G is related to the dipole moment and AG:(i), is the standard free energy of transfer of species i from water into the mixtures, excluding the Born and Kirkwood effects represented by the terms in the brackets on the right-hand side.Subscripts s and w indicate the mixture and pure water, respectively. For the solvolysis2 of 1 ,6-[Copy4C1,]+ in mixtures of water with methanol, propan-2-01 and t-butyl alcohol, eqn (2) must hold: (2) AG:(CZc), # AG:(MZ~), + AG,"(XZx),. 21452146 SOLVOLYSIS OF ~YU~S-DICHLOROTETRAPYRIDINECOBALT(III) Table 1. Mole fractions at which extrema occur in the physical properties of water + cosolvent mixtures and in the transition-state parameters for the solvolysis of 1,6-[COPY,C1,1+ ultrasonic absorption cosolvent (K - c)mina maximumb AHEinc Kmind qmaxe APmaxf AS:axf methanol 0.10 none 0.35 0.15 0.3 0.2 0.15 propan-2-01 0.04 0.15 0.10 0.10 - 0.1 0.1 t-butyl alcohol 0.025 0.10 0.10 0.04 0.4 0.06 0.06 DMSO 0.2-4.39 0.13 0.5 0.4 0.3 - - a Ref.(lo)-( 12). Ref. (1 3) and (14). Ref. (1 l), ( I 5) and (1 6). Ref. (1 7). Ref, (6), (1 1) and (18). f Ref. (2). Q Point of inflection. The physical properties of the mixtures of these alcoholic cosolvents with water,6 particularly the excess change in the temperature of maximum den~ity,~ show that small additions of cosolvent enhance the structure present in liquid water.8 This effect produces a maximum or minimum in some physical properties, first when the maximum stress is exerted within the volumes of structured waters by the packing of the hydrocarbon tails in the spaces outside these volumes, and secondly when the onset of the breakdown of the water structure commences.The first effect is thoughtg to be related to the minimum in the relative partial molar volume of the cosolvent,lo K-- c, whilst the second is indicated by the maximum found in the ultrasonic absorption in the mixture.13 For these properties, and also for the excess enthalpy of mixing, AHE, and for the compressibility, K , the depth or height of the extremum become greater and the mole fraction at which it occurs, x,, becomes lower along the series methanol, propan-2-01 and t-butyl alcohol. Table 1 shows that the movement in the extremum in the enthalpy, AH*, and in the entropy, AS*, of activation for the solvolysis of 1 ,6-[Copy4C1,]+ correlates well with the movement of x2 for the extrema in these physical properties of the mixture.However, table 1 shows that, for water + DMSO (dimethyl sulphoxide), x, for the inflection point in and x, for the extremum in AHE and in K lie well outside the range in x, found for the extrema in these properties in water+alcohol mixtures, even though the extrema in the viscosity, q, and in the ultrasonic absorption all lie in the same range. It is therefore of considerable interest to determine whether AH* and AS* for the solvolysis of 1 ,6-[Copy4C1,]+ in water + DMSO have extrema inside or outside the range found for water + alcohols. - EXPERIMENTAL The materials used were as described previously;2 DMSO was AnalaR grade.The procedure for following the rates of solvolysis was the same as that used with the other cosolvents.2 RESULTS AND DISCUSSION DETERMINATION OF RATE CONSTANTS IN WATER + DMSO Plots of log(optica1 density) against time for the solvolysis in water+DMSO mixtures were always linear for extents of reaction exceeding 85% , as found with the other cosolvents.2 For 35, 40, 45 and 49 "C such plots were obtained for 5.49, 10.92,C. N. ELGY AND C. F. WELLS 2147 Table 2. Values for the first-order rate constant k/10-5 s-' for the solvolysis of 1 ,6-[Copy4C1,]+ in water + DMSO mixtures T I T 34.85 35.00 35.00 35.05 35.10 35.10 35.70 35.80 35.80 37.50 37.50 40.00 40.00 40.05 40.05 40.10 40.70 44.80 44.90 44.90 44.95 45.00 45.05 45.10 47.70 48.00 48.30 48.40 48.45 48.60 48.65 48.70 48.80 48.90 48.95 48.95 49.00 49.00 DMSO (wt%) 5.49 10.92 21.62 32.11 42.38 52.46 2.53 2.92 3.15 - - - 3.30 - - 4.22 4.48 6.4 6.4 6.7 6.5 10.3 11.0 11.3 11.8 16.5 - - - - - - - 18.5 - - - - - - 18.8 17.7 - - 21.62, 32.1 1, 42.38 and 52.46 wt% DMSO, and the first-order rate constants, k, obtained from the slopes are collected in table 2.As there were minor variations in temperature from run to run, the exact temperature is recorded in every case in table 2. Two rate determinations were also made at 37.5 "C for 21.62 wt% DMSO. Good linear plots are obtained when log k is plotted against the reciprocal of the absolute temperature for the same concentration of DMSO. Values for the enthalpy, AH*, and entropy, AS*, of activation were, however, determined by the application2148 SOLVOLYSIS OF trans-DICHLOROTETRAPYRIDINECOBALT(III) Table 3.Values for the enthalpy and entropy of activation for the solvolysis of 1 ,6-[Copy,CI2]+ in water + DMSO mixtures mole fraction AH*/ AS*/ of DMSO kJ mo1-l J K-l mo1-I 0 101 f 11" -3k21" 0.013 109+7 22+ 19 0.027 111+2 2 6 f 6 0.060 108f6 17f 18 0.097 116+4 46+ 12 0.145 122+5 65+ 17 0.202 109f7 26 + 22 1.0 y 0.9 M d + Q 0.8 0.7 a Taken from ref. (2). I I I I I 15 1.27 1.29 1.31 1.3 3 1.35 D , Fig. 1. Plot of log(rate constant) against the reciprocal of the dielectric constant for the solvolysis of 1 ,6-[Copy,C12]+ in water + DMSO mixtures at 25 "C. of the least-squares procedure to all the values of k at their accurate temperatures. These values for AH* and AS* are collected in table 3.INFLUENCE OF SOLVENT STRUCTURE ON THE SOLVOLYSIS Values for the first-order rate constant, k , can be calculated for 25 "C from the values of AH* and AS* in table 3, and values for the dielectric constant D, for water + DMSO mixtures are available for 25 OC.I9 However, although fig. 1 shows that a plot of log k against D;l is approximately linear, this cannot be taken necessarily to imply that solvent structural effects are not involved in the solvolysis of 1,6-[Copy,Cl,]+, since similar plots with other cosolvents added to water are all curved.2 Either the values of AG,O(i), in eqn (1) are all small relative to the first term on the right-hand side, or eqn (3) holds approximately in water + DMSO mixtures: (3) AGr(CZc), z AG:(MZ~), + AG:(XZx)C. N.ELGY AND C. F. WELLS 2149 Fig. 2. Plot of the enthalpy of activation for the solvolysis of 1,6-[Copy,C12]+ in water + DMSO mixtures against mole fraction of DMSO. I0I 70 i Fig. 3. Plot of the entropy of activation at 25 "C for the solvolysis of 1,6-[Copy,C12]+ in water + DMSO mixtures against mole fraction of DMSO. but not for the mixtures of water with the other cosolvents. As curves are obtained with the other cosolvents, implying that eqn (2) holds in these cases with significant magnitudes for AGY(i)n, it would seem unlikely that AG;(i)n (water + DMSO) + AG,"(i) (water + alcoholic cosolvent) and that the second alternative, with eqn (3), is the more likely one. The values for AH* and AS* in table 3 are plotted in fig. 2 and 3 against mole fraction of DMSO in the mixture.Although the variations in AH* and AS* for water + DMSO, relative to the possible deviations, are not so great as those found for the alcoholic cosolvents,2 the trend in AH* and AS* is to continue rising through the composition range where the extrema occur with the branched-chain alcohols as cosolvents, with the subsequent drop for DMSO occurring at about the same x, where2150 SOLVOLYSIS OF ~~~~S-DICHLOROTETRAPYRIDINECOBALT(III) 1 I I 1 1 0 20 40 60 80 J 99' -210 AS*/J K-' mol-I Fig. 4. Plot of the enthalpy of activation against the entropy of activation for the solvolysis of the 1 ,6-[Copy4Cl,]+ ion in water + DMSO mixtures. the extrema occur with methanol. Note that 6 - for water+DMSO mixtures (table 1) shows a point of inflection at about the same x, value at which the falls in AH* and AS* occur.Thus, a water+DMSO mixture does not differ markedly in this respect from the alcoholic cosolvents for the solvolysis of 1,6-[Copy4C1,]+, and the position of the maximum in AH* and AS* in water + DMSO corresponds to the general movement of the physical changes related to solvent structure to higher values of x2 with DMSO compared with the alcohols. Fig. 4 shows that, as with the other cosolvents, the plot of AH* against AS* is linear: a positive deviation on AH* is matched with a positive deviation in AS* and a negative with a negative. For the solvolysis of 1 ,6-[CopyZC1,]+ in water with added methanol, propan-2-01 or t-butyl alcohol,2 as well as for the solvolysis of 1,6-[C0en,Cl,]+,~~ 1,2-[C0en,N,Cl]+,~~ 1,6-[C0en,N,Cl]+,~~ 1 ,6-[Co(4Mepy),Cl2]+ 23 and 1 , ~ - [ C O ~ ~ , S C N C ~ ] + ~ ~ in water + propan-2-01(4Mepy = 4-methylpyridine) and for the ~olvolysis~~ of 176-[Co(4Mepy),- ClJ+ in water + t-butyl alcohol, it has been shown that the influence of the changes in solvent structure on the solvolysis dominates on the pentacoordinated cobalt(II1) ion in the transition state over that on the hexacoordinated cobalt(1Ir) ion in the initial state.This analysis involves the applicationz6 of a free-energy cycle to the solvolysis in water and in the mixture and involves the use of total free energies of transfer of individual species between water and the mixture, AG:(i). For this complex the free-energy cycle will be as follows: A C; [COPY 4c121: * [Copy4Cl]? + c1,C.N. ELGY AND C. F. WELLS 2151 I I I 1 I 0.05 0.10 0.1 5 0.20 0.25 -1 5 XP Fig. 5. Plot of the left-hand side of eqn (4) for the solvolysis of the 1,6-[Copy,Cl,]+ ion in water + DMSO mixtures at 25 "C against mole fraction of DMSO. where w and s indicate pure water and the mixture, respectively. From this cycle at 25 "C the following can be deduced : 2.303RT log (k,/k,) - AGF(Cl-) = AG;(Copy,Cl2+) - AG~(Copy,Cl~) (4) and the experimental values for k , and k, can be combined with the values for AG;(Cl-) derived using the spectrophotometric solvent sorting method.27* 28 It has been found26 for the solvolysis of a wide range of complexes of cobalt(m), as well as for 1 ,6-[Copy4C1,]+ in water + cosolvent mixtures, that the left-hand side of eqn (4) is usually negative.As values for AG;(i) where i is a cation are usually negative, this result implies that the influence of solvent structure dominates in all these cases on the cationic species in the transition state. The above application of the free-energy cycle and of eqn (4) assumes that the Co-..C1 bond is completely broken in the transition state. The evidence in support of the correctness of the assumption for such complexes of ColI1 having two X- ligands and the remaining ligands complexed through four nitrogen atoms, [CoN,X,]+, is good. First, linear plots of log k against the Grunwald-Winstein Y factor29 are usually found,20-22p24 and have always been found2 for [Copy,Cl,]+ in water + cosolvent mixtures: a complete analysis of these plots shows2, why this linearity supports the view that the Co..-Cl bond is as nearly fully extended in the transition state as is the C...Cl bond in the solvolysis of t-butyl Secondly, the magnitudes of volumes of activation, A V*, found for complexes of this type, and particularly the comparison of A V* with the overall volume change, A Vo, indicate full dissociation :31 this has been discussed elsewhere.,, Thirdly, the constancy of the ratio of stereochemical forms found in the product from the solvolysis of [CoN,LX]"+ with varying X for any particular L indicates a common pentacoordinated cation in the transition Fig.5 shows that a plot of the left-hand side of eqn (4) is also negative in water + DMSO for the solvolysis of 1 ,6-[Copy4C1,]+. The values of k , at 25 "C were those determined previously2 and the values for k, at 25 "C were calculated from AH* and AS* in table 3.The values used for AG,"(CI-) in water+ DMSO mixtures at 25 "C had all corrections applied.28 As indicated above, the negative values for the left-hand side of eqn (4) found in fig. 5 show that the effect of change in solvent structure on and the other complexes quoted21 52 SOLVOLYSIS OF ~~-U~S-DICHLOROTETRAPYRIDINECOBALT(III) the cation dominates on the pentacoordinated species in the transition state over the hexacoordinated species in the initial state. The conclusion from this that - AG,"(CopyC12+) > - AG:(Copy,Cl;) in water + DMSO is in accord with similar observations on this complex2 and related complexes20--261 33 in mixtures of cosolvents with water: it is also in accord with -AG:(M2+) > -AG:(M+) usually found for simple cations M of similar size in such mixtures.27* 33 CONCLUSIONS The kinetic results obtained for the solvolysis of 1 ,6-[Copy4C1,]+ in water + DMSO mixtures are in good agreement with those found for the solvolysis of this complex in mixtures of water with alcoholic cosolvents.The extremum in AH* and in AS* moves to higher values of x, than those obtained with the structure-inducing propan-2-01 and t-butyl alcohol, in accordance with the shift in the extremum of physical properties of the mixture related to structure to higher values of x,. The analysis of the free energies of activation shows that this effect of changing solvent structure is greater on the cation in the transition state than in the initial state.Although the approximately linear plot of log k against D;' for x, < 0.2 contrasts with non-linear plots found for the alcoholic cosolvents with this complex, this difference can be compared with the absence of any sharp change in E-- q with DMSO until x, = 0.2 is reached, when an extremum occurs for with alcoholic cosolvents at x, < 0.2. - V. D. Panasyuk and A. V. Arkharov, Russ. J. Inorg. Chem., 1965, 10, 852. C. N. Elgy and C. F. Wells, J . Chern. SOC., Dalton Trans., 1980, 2405; 1982, 1617; J . Chem. SOC., Faraday Trans. I, 1983, 79, 2367. K. J. Laidler and H. Eyring, Ann. N. Y. Acad. Sci., 1940, 39, 303; S. Glasstone, K. J. Laidler and H. Eyring, The Theory of Rate Processes (McGraw-Hill, New York, 1941), chap. 8; K. J. Laidler and P.A. Landskroener, Trans. Faraday Soc., 1956, 52, 200; K. J. Laidler, Suom. Kemistil. A, 1960, 33, 44; K. J. Laidler, Chemical Kinetics (McGraw-Hill, New York, 2nd edn, 1965), chap. 5. E. A. Meolwyn-Hughes, Proc. R. Soc. London, Ser. A, 1936, 155, 308; 1936, 157, 667; The Kinetics of Reactions in Solution (Oxford University Press, London, 2nd edn, 1947), chap. 4, 5 and 7; Physical Chemistry (Pergamon Press, Oxford, 2nd edn, 1961), chap. 7-9 and 24. E. S. Amis, Kinetics of Chemical Change in Solution (Macmillan, New York, 1949), chap. 5 and 9; Solvent Eflects on Reaction Rates and Mechanisms (Academic Press, New York, 1966), chap. 1-3; Solvent Eflects on Chemical Phenomena (Academic Press, New York, 1973), vol. 1, chap. 5. F. Franks and D. J. G. Ives, Quart.Rev. Chem. SOC., 1966, 20, 1. G. Wada and S. Umeda, Buff. Chem. SOC. Jpn, 1962, 35, 646. H. S. Frank and M. W. Evans, J. Chem. Phys., 1945, 13, 507; H. S. Frank and W-Y. Wen, Discuss. Faraday SOC., 1957, 24, 133; G. Nemethy and H. A. Sheraga, J. Chem. Phys., 1962, 36, 3382, 3401; N. Laiden and G. Nemethy, J. Phys. Chem., 1970,74, 3501. C. F. Wells, Trans. Faraday Soc., 1970, 66, 204. l o K. Nakanishi, Buff. Chem. SOC. Jpn, 1960, 33, 793. l 1 J. Kenttamaa, E. Tommila and M. Martti, Ann. Acad. Scient. Fenn., Ser. A, 1959, no. 93. l2 J. Kenttamaa and J. J. Lindberg, Suom. Kemistif. B, 1960, 33, 32. l3 M. J. Blandamer, Introduction to Chemical Ultrasonics (Academic Press, London, 1973), chap. 1 1. l4 D. E. Bowen, M. A. Priesand and M. P. Eastman, J. Phys. Chem., 1974, 78, 261 1.l5 R. F. Lama and B. C-Y. Lu, J . Chem. Eng. Data, 1965, 10, 216. l6 J. Kenttamaa and J. J. Lindberg, Suom. Kemistif. B, 1960, 33, 98. l i K. H. Jung and J. B. Hyne, Can. J . Chem., 1970,48, 2423. R. G. LeBel and D. A. I. Goring J. Chem. Eng. Data, 1962, 7, 100. J. J. Lindberg and J. Kenttamaa, Suom. Kemistil. B, 1960, 33, 104. 2o G. S. Groves and C. F. Wells, J. Chem. Soc., Faraday Trans. I , 1983, 79, 253. *l A. E. Eid and C. F. Wells, J. Chem. Soc., Faraday Trans. I, 1981, 77, 1621. 22 A. E. Eid and C. F. Wells, J. Chem. Soc., Faraday Trans. I , 1983, 79, 253. 23 I. M. Sidahmed and C. F. Wells, J. Chem. SOC., Dalton Trans., 1983, 1035.C . N. ELGY AND C . F. WELLS 2153 24 A. E. Eid and C. F. Wells, J . Chem. SOC., Faraday Trans. I , 1985, 81, 1401. 25 I. M. Sidahmed and C. F. Wells, J . Chem. SOC., Dalton Trans., 1984, 1969. 26 C. F. Wells, J. Chem. Soc., Faraday Trans. I, 1977, 73, 1851. 27 C. F. Wells, J. Chem. SOC., Faraday Trans. I , 1973, 69, 984; 1974, 70, 694; 1975, 71, 1868; 1976,72, 601; 1978, 74, 636, 1569; 1981, 77, 1515; 1984, 80. 2445; G. S. Groves, I. M. Sidahmed and C. F. Wells, unpublished work. 28 C. F. Wells, Ausrr. J. Chem., 1983, 36, 1739. 89 E. Grunwald and S. Winstein, J. Am. Chem. SOC., 1948, 70, 846. 30 M. H. Abraham, Prog. Phys. Org. Chem., 1974, 11, 1. W. E. Jones, L. R. Carey and J. W. Swaddle, Can. J. Chem., 1972,50, 2739; G. A. Lawrance, Inorg. Chim. Acta, 1980, 45, L275; G. A. Lawrance and S. Suvachittanont, Austr. J. Chem., 1980, 33, 277; D. A. Palmer and H. Kelm, Inorg. Chem., 1977,16,3 139; Coord. Chem. Rev., 198 I,%, 89; G. Daffner, D. A. Palmer and H. Kelm, Inorg. Chim. Acta, 1982, 61, 57. 32 W. G. Jackson and A. M. Sargeson, Inorg. Chem., 1978,17, 1348; W. G. Jackson and C. M. Begbie, Inorg. Chim. Acta, 1982, 60, 115. 33 I. M. Sidahmed and C. F. Wells, J. Chem. SOC.. Dalton Trans., 1981, 2034. (PAPER 4/ 198 1 )

ISSN:0300-9599    

DOI:10.1039/F19858102145

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. I , 1985, 81, 2155-2168 Interfacial Tension Minima in Oil +Water + Surfactant Systems Effects of Salt and Temperature in Systems containing Non-ionic Surfactants BY ROBERT AVEYARD,* BERNARD P. BINKS, THOMAS A: LAWLESS AND JEREMY MEAD Department of Chemistry, University of Hull, Hull HU6 7RX Receiued 26th November, 1984 Oil/water interfacial tensions have been determined in systems containing alkane, aqueous NaCl and N-dodecyl polyoxyethylene non-ionic surfactants. The tensions are frequently very low (ca. mN m-l) and pass through a minimum as the salt concentration, temperature or alkane chain length is varied. By the use of thermodynamics it is shown that the minimum with respect to salt concentration can arise through the competing effects on tension of the positive adsorption of surfactant and the negative adsorption of salt at the oil/water interface.It is assumed that surfactant adsorption is independent of salt concentration; it is not necessary to suppose that the sign of the surface excess of surfactant changes as a tension minimum is traversed. It is shown that a minimum in tension with respect to temperature can occur when the entropy of micelle formation and of formation of the plane oil/water interface (both expressed per mole of surfactant) are equal. There is a continuing interest in microemulsion systems containing (at least) oil, water and either ionic or non-ionic surfactant~.l-~ In both two- and three-phase regimes in such systems very low interfacial tensions can be attained. Our interest here centres on these tensions and in particular on the way in which they vary with inorganic salt concentration and temperature.We have obtained a large body of experimental data on chemically pure systems for both ionic and non-ionic surfactants and much of this will be reported elsewhere. The purpose of the present paper is to explore the significance of oil/water interfacial tension minima in systems containing alkane, aqueous NaCl and pure polyoxyethylene surfactants of the type n- C,,H,,(OCH,CH,),OH(designated C,,E,). Recently Ruckenstein and Beunen5 have proposed the possibility that tension minima observed with respect to salt concentration (at high salt concentration) can arise as a result of the change in sign of the surface excess of surfactant at the oil/water interface.This could mean that at high salt concentrations (say around 1 mol dm-3 NaCl), when the interfacial tension is still very low (say - 10-l mN m-I) the surface excess of surfactant is negative. Here we give a thermodynamic treatment leading to an alternative and, we believe, more likely explanation of tension minima in systems containing non-ionic surfactant and 1,l -electrolyte. We also consider minima in tension with respect to temperature. EXPERIMENTAL High interfacial tensions (i.e. ca. 50 mN m-l) were measured using the drop-volume technique; tensions were reproducible to better 0.05 mN m-1.6 The low interfacial tensions (< 2 mN m-l) were measured using a Kruss spinning-drop tensiometer. A small drop (ca. 21552156 TENSION MINIMA IN NON-IONIC SURFACTANT SYSTEMS O! 0 4000 60 n/Hz 8 000 Fig.1. Plot of the recorded tension, y, against the frequency of rotation n of the capillary for the system C,,E, + nonane + aqueous NaCl (1.35 mol dmP3) at 31 "C. 1 mm3, depending on the tension) of oil is injected into the aqueous phase in a horizontal capillary tube rotating about its long axis. At sufficiently high frequencies the drop assumes the form of a cylinder, radius r , which hemispherical ends. Under these conditions the interfacial tension y is given by y = 0.25r3Apco2 (1) where o is the angular velocity (i.e. 2nn, where n is the frequency of rotation) and Ap is the difference in density between the two phases. For the most part tensions reached a steady mean value within a few minutes and remained unchanged over a period of at least 4 h, although some periodic oscillations were observed, as described later.Recorded tensions were independent of the frequency of rotation, as can be seen from fig. 1. Phase inversion from oil-in-water (o/w) to water-in-oil (w/o) emulsions was detected by determination of the conductance of the emulsion using a digital conductivity meter with an added facility for the digital read-out of temperature (Jenway, model PCM3). The distribution of surfactants between alkane and aqueous phases was determined as a function of surfactant and salt concentrations. Equal volumes of heptane and pure water were shaken with various amounts of C,,E, and the resulting emulsions were allowed to separate in a water thermostat at 25 "C for 48 h.The aqueous and alkane phases were then centrifuged at 25 "C for 1 h before analysis. For study of the effects of salt concentration, one volume of 0.043 mol dm-3 C,,E, in nonane was shaken with five volumes of aqueous solution containing various concentrations of NaCl. The phases were equilibrated and separated as previously but at 3 1 "C. The surfactant concentrations in each phase were determined spectrophotometrically by complexation with ammonium cobaltothiocyanate; the complex absorbs strongly at 325 nm in benzene.' Photon correlation spectroscopy (P.c.s.) measurements on the separate phases were recorded using a Malvern Instruments PCS 100 spectrometer. Polarised light from a Spectra-Physics 15 mW He-Ne laser was focussed on the sample contained in a 1 cm fluorescence cuvette suspended in a dish containing thermostatted water.The scattered light was detected at 90" to the incident beam and the correlation function evaluated by a Malvern K7027 correlator. The correlation functions were analysed by the method of cumulants to give a mean value of the correlation length and a variance of the distribution of exponentials.* The correlation lengthR . AVEYARD, B. P. BINKS, T. A. LAWLESS AND J . MEAD 2157 5 15 T/"C 25 Fig. 2. Effect of initial oil-phase surfactant concentration on tensions for the system C,,E, + octane+0.085 mol dm-3 NaCl. Oil-phase concentrations represented are (in wt% ) 5.0, 3.8,2.7 and 1.8. At the lowest temperatures, the higher concentrations give the lower tensions; otherwise tensions are independent of concentration.may be related to the hydrodynamic radius in the case of non-interacting dispersions (or in the limit of infinite dilution). The surfactants used were pure homogeneous samples supplied by the Nikko Chemicals Co. (Japan) and were used as received. Water was distilled once, passed through an Elgastat ion-exchange column and then through a Milli-Q reagent water system. Water treated in this way had a surface tension against air (determined by various methods) in excellent agreement with the best literature values. Alkanes were from various commercial sources and were > 99% pure as supplied; g.1.c. analyses confirmed these purities. Samples were passed through chromatographic alumina before use to remove traces of polar impurities.Sodium chloride was AnalaR grade and was used as received. It was confirmed, however, that aqueous solutions had surface tensions in good agreement with those previously obtained using samples of NaCl roasted for 12 h at 450 O C 6 RESULTS AND DISCUSSION SOME CHARACTERISTICS OF THE SYSTEMS In work similar to that described here but using anionic surfactants we have found that ultralow oil/water tensions can be obtained by injecting a drop of pure alkane into an aqueous surfactant at or above its critical micelle concentration (c.m.c). If this procedure is used for C,,E, surfactants, however, ultralow y do not result and we have21 58 TENSION MINIMA IN NON-IONIC SURFACTANT SYSTEMS 0 1 2 6 8 10 12 14 16 [NaCl]/mol drn+ N Fig. 3. (a) Plot of yc against salt concentration for C,,E,+nonane+NaCl at 31 "C.(b) Plot of yc against alkane chain length for C,,E, at 40.0 "C. found it necessary to put the surfactant not in the aqueous phase but the oil phase and at a high concentration (say 0.02 mol dm-3) relative to the aqueous phase c.m.c. Above a certain minimum concentration, however, y is found to be independent of the concentration (fig. 2). We have also determined tensions in oil + water systems that have been allowed to come to equilibrium prior to being introduced into the spinning- drop apparatus. Tensions so obtained were in excellent agreement with those for systems not originally at equilibrium. In fig. 2 and 3 we show how y varies with temperature, aqueous-phase NaCl concentration and alkane chain length, N .In all cases y passes through a sharp minimum. In the y against T and y against N curves the aqueous phase was 0.085 mol dm-3 NaCl. Such a low concentration has little effect on y (as seen from the y against [salt] curve) and was used in anticipation of work to be done on non-ionic and anionic surfactant mixtures where the presence of NaCl is required. In order to understand why relatively high oil-phase concentrations of surfactant are required to obtain low y we have investigated the equilibrium distribution of surfactant between alkane and water as a function of both surfactant and salt concentrations; our findings are summarised in fig. 4. Results for the distribution of C,,E, between heptane and water at 25 "C are shown in fig. 4(a), where it is seen that as the initial oil-phase surfactant concentration is increased, the equilibrium oil-phase concentration rises and the aqueous-phase concentration, which is very low, changes only little.As the initial surfactant concentration is increased further the equilibrium oil-phase concentration reaches a plateau and the aqueous-phase concentration begins to rise. The effect of salt concentration on the distribution of (a fixed total amount of) C,,E, between water and nonane at 31 "C is shown in fig. 4(b). At low salt concentrations the aqueous- and alkane-phase concentrations are high, but as the salt concentration increases the aqueous-phase concentration falls to a low and almostR. AVEYARD, B. P. BINKS, T. A. LAWLESS AND J. MEAD 2159 E E 2 6 c) 0 2 L 6 [initial surfactant] in heptane/ mol dm-3 15 O 1 a E N N E 0 0.5 1.0 1.5 [salt]/mol dm-3 Fig. 4.Distribution of surfactant between phases. Filled circles: surfactant in water; open circles: surfactant in alkane. (a) Effect of surfactant concentration for C,,E, + heptane + water at 25 "C. (b) Effect of salt concentration for C,,E,+nonane+NaCl at 31 "C. constant value, approximately equal to reported c.m.c. values, and the oil-phase concentration rises to a high plateau value. We have interpreted these findings as follows. Monomeric surfactant distributes strongly in favour of alkane. This is to be expected and is consistent with the distribution data of Manabe et al.,9 who studied similar materials but with shorter alkyl chains. As the concentration of the surfactant in the system increases an aggregation point (which we term the c.m.c.) is reached either in the aqueous phase or in the oil phase, depending on conditions (T, salt concentration and N).As a result of the way in which monomer distributes, the c.m.c. condition corresponds to a high concentration in the oil phase and a low concentration in the aqueous phase. Above the c.m.c. (at constant 7' and salt concentration) the monomer activity remains essentially constant, and since it is the monomer that is surface active the tension also remains constant. For the system represented in fig. 4(a) the aggregates form in the2160 1 0 TENSION MINIMA IN NON-IONIC SURFACTANT SYSTEMS 50 100 delay time/ps 150 Fig. 5. Semi-logarithmic plot of the correlation function, G(t), against the delay time for an equilibrated aqueous phase in the system C,,E, + heptane + water at 25 "C.The solid line shows the mean decay rate derived from the cumulants analysis. aqueous phase rather than the oil phase, and so the aqueous-phase concentration rises with overall surfactant concentration. In fig. 4(6) the concentration of surfactant in the system is constant. At low salt concentrations the surfactant concentration in the aqueous phase is high and an oil-in-water microemulsion is present. As the salt concentration is increased surfactant transfers to the oil phase and forms aggregates (in the form of water-in-oil microemulsion droplets coated with a surfactant monolayer) in equilibrium with monomers at highconcentration. The aqueous-phase concentration falls to a level close to the c.m.c.that would be observed in the absence of excess oil phase. The above propositions are given more weight by photon correlation spectroscopy data obtained for the equilibrium phases. First we consider the variation of the surfactant concentration in pure water. None of the heptane phases exhibited any correlation function over a range of delay times appropriate for aggregate radii from < 1 to 250 nm. This was also true of the three aqueous phases with the lowest surfactant concentrations [fig. 4(a)]. For the other (seven) aqueous phases P.C.S. showed that small aggregates were present in solution and the correlation lengths increased with concentration in the range 8-1 1 nm. Variances were all in the range 5-10% and a typical plot of the logarithm of the correlation function against delay time is shown in fig.5 . Extrapolation of the correlation lengths to infinite dilution (i.e. assuming the size of the aggregates to be independent of concentration) gives a hydrodynamic radius of ca. 7.5 nm. Some preliminary experiments were also performed on aqueous and alkane phases either side of the salt concentration causing phase inversion. At low salt concentrations structure was found entirely in the aqueous phase. At high salt concentrations, however, no structure was found in the aqueous phase but the nonane phase exhibited light scattering typical of small water-in-oil microemulsion droplets.R. AVEYARD, B. P. BINKS, T. A. LAWLESS AND J . MEAD 2161 20 30 40 T I T 2000 1500 4 I 9 1000 $ 0 Q s 500 0 Fig.6. Plots of conductances and interfacial tensions against temperature for the system C,,E, + hexane +0.085 mol dm-3 NaC1. Open circles are conductances and filledcircles tensions. Curves (a), (b) and (c) are for phase-volume ratios (water:oil) of, respectively, 2: 1, 1 : 1 and 1 : 2. Conductances for the 2: 1 data have been divided by 1.8 to bring them on scale. From the foregoing considerations we may suppose that as surfactant concentration in a system is increased, y falls until the onset of aggregation in the preferred phase and then remains constant at a value of yc. This kind of tension variation has been observed by Crook et aZ.l0 for octylphenoxyethoxyethanols (OPE) with E, to El, groups in water + isooctane systems. Thus, the y against [salt], y against Tand y against N curves are seen to be plots of yc against the appropriate variable.The conditions corresponding to minimum yc with respect to [salt], T and N also correspond to phase inversion, i.e. the point where emulsions formed from the oil + water systems change from the oil-in-water (o/w) to the water-in-oil (w/o) type. This is illustrated in fig. 6. At low temperatures, where surfactant forms micelles in the aqueous phase, o/w emulsions are formed and exhibit high conductivity. Above T* (T corresponding to minimum ye) surfactant resides mainly in the alkane phase (see later) and w/o emulsions result. Here the continuous phase is alkane and the conductivity is correspondingly low. Phase inversion and surfactant transfer between phases is probably associated with the effective geometry of the surfactant molecules.2 The latter determines the preferred curvature of the interface containing the surfactant.For swollen micelles in water the effective cross-sectional area of the head group exceeds that of the chain, and for water droplets in w/o microemulsions the chain area is greater than that of the head group. At phase inversion (and minimum ye) the two areas are equal and zero surface curvature is the preferred state. Note that T* is 2 or 3 "C higher than the cloud point of the aqueous phase saturated with alkane, 71 F A R 812162 TENSION MINIMA IN NON-IONIC SURFACTANT SYSTEMS but in the absence of an excess of the oil phase. Thus, in simple terms, if surfactant is not able to transfer to an oil phase, it will instead phase separate.The relationship between cloud points, formation of a third phase (see below) and phase inversion will be discussed in a later publication. Before proceeding to a thermodynamic treatment of the occurrence of tension minima, we mention two interesting observations made concerning tension measure- ments. During measurements leading to the yc against Tcurve it was noticed that as the temperature was lowered, just above T* a third phase was formed, which appeared initially as a coating on the oil drop and then, at lower T, as segments in between segments of the oil drop. A sequence of photographs of a drop taken as the temperature is lowered is shown in plate 1. At the higher temperatures surfactant is present largely in the oil phase, then largely in the third phase, which, as Tis lowered further, is dispersed in the aqueous phase, as is evident from plate 1. The observation of the formation of a third phase in such systems is entirely in accord with results of various phase-equilibria studies reported in the literature [e.g.ref. (4) and (1 l)]. The second observation is that in both pre-equilibrated and non-pre-equilibrated systems it was often noted that the tension varied in an oscillatory fashion with a period of ca. 1 min; the oscillations persisted, about the same mean y, over long periods of time. The effect, which is not observed for anionic surfactants, could arise as a result of the slow rates of diffusion of the non-ionic surfactants to and from the interface. When anoil drop is introduced into the capillary of the spinning-drop apparatus, it elongates and its surface area increases, leading, if the surface concentration of surfactant is not maintained, to an increase in y.This in turn would tend to cause a contraction of the drop, giving an increase in surface concentration and a concommitant fall in y, followed by elongation of the drop and so on. THERMODYNAMIC TREATMENT VARIATION OF yc WITH m, We consider a system comprising a water-insoluble oil (alkane say), an aqueous solution containing a 1 : 1 electrolyte (which we designate S) and a non-ionic surfactant D distributed between the two phases as described. We employ the following terms: y is the tension between the oil and aqueous phases and yc refers to systems at and above the c.m.c.(as defined earlier), mD, m,, m+ and m- are the molar concentrations in aqueous phase of D, S, cation and anion, respectively, TD, Tcl-, rNa+ and T, are the surface excesses of surfactant, anion, cation and salt, respectively, f+,f- and f+ are the molar activity coefficients of cation, anion and mean ionic activity coefficient of S, respectively, a+ and a- are the molar activities of cation and anion, respectively (a = m f ) andf, is the molar activity coefficient of D in aqueous solution. On the basis of the model, y on the y against [salt], y against T and y against N curves (fig. 2 and 3) are all yc. Consider a system at constant Tin which the y against In m, curves, each for fixed m,, are as depicted in fig. 7. The tension difference between points A and B, which are at the c.m.c.for salt concentrations ms,l and m,,,, respectively, is given by the sum of the differences in tension between points A and C and between points C and B, i.e. The terms a y / a In mD and a y / a In rn, can be obtained from the Gibbs equation, which for constant pressure is -dy = C Tidpu,+S<dT. (3) zJ . Chem. SOC., Faraday Trans. 1, Vol. 81, part 9 Plate 1 AVEYARD, BINKS, LAWLESS AND MEAD (Facing p . 2 162)R. AVEYARD, B. P. BINKS, T. A. LAWLESS AND J. MEAD 2163 B In m,, Fig. 7. Schematic representation of y against In rn, for two salt concentrations. We will for the moment employ the Gibbs convention for the interface and choose a dividing surface such that the surface excess of water is zero; then the Ti are relative adsorptions and is the relative surface entropy.12 Furthermore, at the c.m.c.the mole fraction of surfactant in the oil phase is only of the order of 5 x at most, so we may reasonably take dp (oil) to be negligible. Thus, the species to be included in eqn (3) (neglecting H+ and OH-) are Na+, C1- and D, so that at constant T - dy = RT(T,,+ d In a+ + Tcl - d In acl- + r D d In a,,). (4) From eqn (4), noting that for electroneutrality of the interface rNa+ = Tcl- = Ts, we Furthermore, by assuming that low concentrations of D (i.e. mD < c.m.c.) do not significantly affect f+ and that fD does not vary with mD (i.e. the surfactant is in infinitely dilute idealsolution) Combination of eqn (2), (5) and (6), noting that in the case of interest y = yc and m, = c.m.c., yields -- - a lnrn - 2.303 RTrD [ D drnS The quantity d logfD/drns is frequently constant up to rn, = 1 or 2 mol dm-3 and beyond and is the salting constant k, for the surfactant monomer in s01ution.l~ Positive values correspond to salting-out and negative values to salting-in.The quantity k, can usefully be split into additive group contributions for the non-polar chain, k,,, and for the polar head group, kp.14 The term in the c.m.c. in eqn (7) is also known to be constant for a number of non-ionic s ~ r f a c t a n t s ~ ~ and is equal to 71-22164 TENSION MINIMA IN NON-IONIC SURFACTANT SYSTEMS E z 1 E --. 0 0 .I 0.2 0.3 C12 E, in heptane (10' mole fraction) Fig. 8. Surface pressure, n, of C,,E, adsorbed from heptane to the heptane/aqueous NaCl interface at 25 "C.Points 0, a, and 0 refer to ms = 0, 1.02 and 2.06 rnol dm-3, respectively. km-k,, where k , is the salting constant for the monomer in the micelle.16 The quantity k , is frequently assumed to arise mainly from the effect of the salt on the polar group, i.e. k, % kp, ,. Inspection of eqn (7) reveals the possibility of a minimum in yc with respect to m, without TD changing sign. For a salted-out monomer such that k, > Id log c.m.c./ dm, 1, the term in TD is positive for positive TD. Sodium chloride (for example) is likely to be desorbed at the oil/water interface6 so the term in Ts will be negative since 13 lnf+/a lnm, I < 1. The magnitude of the positive term in T D can vary through variatrons in T,, with m,, and the negative term in Ts can vary through changes in Ts/m, and the term inf+.Clearly then the sign of dy,/dm, can change and yc pass through a minimum if the terms in T,, and Ts are finely balanced. We now explore how well eqn (7) is able to describe the curves observed experimentally. First we consider what can be simply deduced from the position of the minimum in the y against [salt] curve. In the system containing CI2E,+nonane [fig. 3(a)] the minimum yc occurs at m, = 0.5 mol dmP3. Suppose TD at m, = 0.5 mol dm-3 is equal to the value for zero salt, i.e. 3.3 x mol m-2;17 that this is reasonable will be seen later. The value of d log c.m.c./dm, will be taken as -0.41 dm3 mol-l, a value obtained for C,,E6 in the absence of alkane [as estimated from fig. 1 of ref. (1 5)]. The value of Ts is taken to be equal to that for 0.5 mol dm-3 NaCl in contact with decanol,l* which is -0.83 x lo-' mol m-2.The interface with decanol may not be too dissimilar to the interface containing a saturated monolayer of non-ionic surfactant. Logarithms off+(NaC1)lg have been fitted to a cubic in In m, to obtain 8 lnf+ /? In m,. At m, = m:, The salt concentration at minimum yc, it can be seen from e<n (7) and the discussion following it that a lnf + s [ l + ( - ) m: alnm, 1. (8) = 2.303 r D k p , , = d log c.m.c. 2.303 TD (k,+ dm, The value of k, obtained by the use of eqn (8) is 0.45 and k p , = 0.04. This indicates (if the k values are reasonable) that the effect which salt has on the c.m.c. is mainlyR. AVEYARD, B. P. BINKS, T. A. LAWLESS AND J. MEAD 2165 a result of the salting-out of the monomer from solution, the salting-out of the micellar surfactant being very small.Nishikido and Matuura15 have obtained a value of k,, = 0.04, k, is thus seen to be - 0.13, i.e. the E group appears to be salted-in in bulk. It has been suggested by Schickz0 that it is the salting-out of (larger) E groups which is responsible for the depression of the c.m.c. by salts. However, Ray and Nemethy21 found that salt effects on the c.m.c. of surfactants containing E, and E,, groups were identical. Furthermore, Mukerjee16 maintains that the salt effects arise in the main from salting-out of the non-polar species in bulk solution. Interestingly, we have determined directly the salt effect on isolated E, groups (as opposed to close-packed groups at a micelle surface) at an oil/water interface and find them to be salted-in by NaCl.In fig. 8 we show surface pressures 71 of C,,E, adsorbed from heptane to the interface with water or salt solution. It is seen that for a given surfactant concentration in heptane, 71 is increased by addition of salt to the aqueous phase, which means surfactant is more strongly adsorbed when the E, is immersed in salt solution at the interface. From these data a surface salting-in coefficient, 6, of -0.08 is obtained.,, In any event, from the various values discussed above we obtain for k,, (for the C,, alkyl chain in aqueous solution) a value of 0.58. In a successful semi-empirical approach to salting-out of non-polar groups and molecules presented by Aveyard,,, k,, is given as - k, of 0.17 for C1,E6 with NaCl.For k,, k,, = (N/2.303 RT) (a,, Ay) (9) where Rnp is the surface area of the non-polar group and Ay is the increase in the surface tension of water caused by addition of 1 mol dm-3 salt. Taking values of R,, = 304A2,, and Ay = 1.7 mN m-1,6 k,, calculated from eqn (9) is 0.54, in close agreement with the value of 0.58 obtained above. We now test eqn (7) to see how well it reproduces a complete yc against [salt] curve for C12E5 with NaCl. Using values of Ts from ref. (18) andf, - from ref. (19), the term in Ts in eqn (7) is closely fitted by 2r, (1 +%) = - 1.065 x m, alnm, exp(-6.829~~~)-3.05 x with Ts in mol m-,. As before d log c.m.c./dm, has been set at - 0.41 l5 and r D is taken as constant and equal to 3.3 x lop6 mol m-,.The salting constant d logfD/dms is taken as an adjustable constant. Eqn (7) in integrated form then becomes y = -RT[2.303TD(k,-0.41)~- 1.065 x exp(-6.829ms)-3.05 x 10-7m,]+B where the integration constant B is yo - 1.065 x 1 O-’ RT, yo being the tension at m, = 0. The fit of the experimental data to eqn (lo) was optimised with respect to i3 logfD/i3mS giving a value of 0.4455 and the curve generated using these values is shown in fig. 9. The general form of the yc against [salt] curve is well reproduced and the minimum occurs at the correct salt concentration; the value of k, obtained from the fit has already been shown to be realistic. Note that the fit is obtained assuming TD is constant with respect to salt concentration. It appears to us on intuitive grounds that even if TD does in fact vary with m, it is very unlikely to become negative.(10) VARIATION OF yc WITH T We turn now to the shape of the yc against T curve and its significance. It will lead to greater clarity here if we adopt the ‘surface-phase’ approach (in which the surface entropy and surface concentrations are total quantities) rather than the Gibbs model2166 TENSION MINIMA IN NON-IONIC SURFACTANT SYSTEMS 0 1 [NaCll/mol dm-3 2 Fig. 9. Plot of the fit of the experimental yc against [salt] curve for C,,E, + nonane + NaCl using eqn (10) with k, = 0.4455 and B = -0.0487 mN m-l. Points are experimental and the line represents the fit. Filled points are for pre-equilibrated phases obtained from distribution experiments. Error bars show oscillatory variation in value of yc.for the surface, in which excess quantities appear. We denote the total quantities by the superscript s. Salt is not necessary to obtain a minimum in yc with respect to T and is not therefore included in the treatment which follows. For changes in yc with T we may write where x represents mole fraction. We denote the oil phase by a and the aqueous phase by p, with subscripts o and w refering to oil (alkane) and water, respectively. We will suppose that x t and x", are effectively zero and that the surfactant D is distributed between a and B phases. The Gibbs equation for this system is (12) -dy = SS, dT+ TS, dp, + T$ dpw + Tb dpD.R. AVEYARD, B. P. BINKS, T. A. LAWLESS AND J. MEAD 2167 where the Si are partial molar entropies and Si is the entropy per unit area of interface.For changes maintaining equilibrium, d& = dpk and so from eqn (15) and (16), assuming ideal behaviour of surfactant below the c.m.c., RTd lnxf, = (Sg-Sf,)dT+RTd lnxf,. (17) Since the solutions of D in both oil and water are very dilute we may reasonably neglect terms in (8p:/C)~%)~ and (8&/ax&),.. Hence, on the assumption of ideal behaviour of surfactant below the c.m.c. (non-ideality, particularly in the oil phase, could be a complicating factor), combination of eqn (1 2)-( 1 5 ) yields -dy = (ss,-r;sz-p sb’-p sa ,)dT+RTrS,dlnxa, ( a ~ p T),: = - (s; - I-; s: - I-; st - r b ss) (18) (19) (Q@ lnxff), = -RTI-S,. (20) so that Noting that -RTd lnc.m.c./dT= AS,, the entropy of micelle formation at the c . ~ . c ., ~ ~ combination of eqn ( 1 l), (19) and (20) for x6 = c.m.c. in a, yields Eqn (21) is appropriate where micelles form in the oil phase, i.e. on the high-T limb of the yc against T curve. Combination of eqn (1 7) and (21) leads to which applies to the low-T limb where micelles form in phase p. The entropy of micelle formation is given by25 AS, = S,-C N i s i i where S , is the entropy of micelles containing one mole of surfactant ( N , = 1 ) together with the other components and the Si are the partial molar entropies in bulk solution of all the i components involved. The entropy terms in parentheses in eqn (21) and (22) are excess surface entropies of unit area of interface, which contains Tb moles of surfactant. It is thus seen that the form of the yc against Tcurve arises from the differences between the entropy of forming a plane interface and the entropy of forming micelles containing the same amount of surfactant.At the temperature P corresponding to the minimum yc (i.e. the phase-inversion temperature), the entropy of micelle formation from the constituents in bulk is equal to the entropy of surface formation, both entropies being expressed per mole of surfactant. On the low-T limb of the yc against Tcurve surfactant is aggregated in the aqueous phase as micelles containing solubilised oil and dy,/dT is negative. Ask is positive for C12E, and becomes less positive with increasing T,15 so for T < T*, l-6 AS& is more positive than the entropy of surface formation appearing in eqn (22). For T > T* the aqueous phase is probably at its c.m.c., with no micelles present and the excess surfactant is in the oil, which is likely to be a dilute w/o microemulsion.In this case the entropy of surface formation in eqn (21) is more positive than rgAS&. If one accepts the ‘geometrical’ picture of phase inversion,2 the surfactant resides in water when the E group is large as a result of hydration and ‘prefers’ to be on the convex side of a curved interface (i.e. the micelle in water). Transfer of monomer to micelle might be expected to cause some dehydration of the E groups (as well as of the alkyl chain)2168 TENSION MINIMA IN NON-IONIC SURFACTANT SYSTEMS in close proximity at the micelle surface. This would give a positive contribution to AS&. As the temperature is increased, the solvation (and hence effective size) of the head group decreases and AS& becomes less positive. When AS& becomes equal to the entropy of surface formation, zero (net) curvature is the preferred state and the head-group and the effective chain cross-sectional areas become equal.Further decrease in AS& implies that the head group prefers to be on the concave side of an interface, i.e. on a w/o microemulsion droplet surface. We thank the S.E.R.C. and British Petroleum for the award of a Cooperative Grant and British Petroleum for EMRA funding and the award of a Research Studentship (to B.P.B.). We also thank Miss H. Slezok for measurements leading to the data presented in fig. 8. J. Th. G. Overbeek, in Surfacfanfs, ed. Th. F. Tadros (Academic Press, London, 1984), chap. 5. D. J. Mitchell and B. W. Ninham, J. Chem. Soc., Faraday Trans. 2, 1981,77,601. A. M. Cazabat, D. Langevin, J. Meunier and A. Pouchelon, Adv. Colloid Interface Sci., 1982, 16, 175. H. Kunieda and K. Shinoda, Bull. Chem. SOC. Jpn, 1982, 55, 1777. E. Ruckenstein and J. A. Beunen, J. Colloid Interface Sci., 1984, 98, 55. R. Aveyard and S. M. Saleem, J. Chem. SOC., Faraday Trans. 1, 1976, 72, 1609. D. E. Koppel, J. Chem. Phys., 1972, 57, 48 14. M. Manabe, M. Koda and K. Shirahasma, Bull. Chem. SOC. Jpn, 1975 48, 3553. ’ R. A. Greff, E. M. Setzkom and W. D. Leslie, J . Am. Oil Chem. Soc., 1965,42, 180. lo E. H. Crook, D. B. Fordyce and G. F. Trebbi, J. Phys. Chem., 1963,67, 1987. l1 H. Kunieda and K. Shinoda, J. Dispersion Sci. Technol., 1982, 3, 233. l 2 R. Aveyard, B. P. Binks and J. Mead, J. Chem. Soc., Faraday Trans. 1, 1985, 81, 2169. l3 F. Long and W. McDevit, Chem. Rev., 1952,51, 1 19. l4 F. L. Wilkox and E. E. Schrier, J. Phys. Chem., 1971,75, 3757. l 5 N. Nishikido and R. Matuura, Bull. Chem. SOC. Jpn, 1977, 50, 1690. l6 P. Mukerjee, J. Phys. Chem., 1965, 69, 4038. l 7 M. J. Rosen, A. W. Cohen, M. Dahanayake and X. Hua, J. Phys. Chem., 1982,86, 541. l9 R. A. Robinson and R. H. Stokes, Electroiyte Solutions (Butterworths, Sevenoaks, 1955). 2o M. J. Schick, J. Colloid Sci., 1962, 17, 801. 21 A. Ray and G. Nemethy, J. Am. Chem. Soc., 1971,93, 6887. 22 R. Aveyard and S. M. Saleem, Can. J. Chem., 1977, 55, 4018. 23 R. Aveyard, Can. J. Chem., 1982, 60, 1317. 24 E. A. Guggenheim, Thermodynamics (North Holland, Amsterdam, 5th edn, 1967). 25 D. G. Hall, in Aggregation Processes in Solution, ed. E. Wyn-Jones and J. Gormally (Elsevier, . R. Aveyard, S. M. Saleem and R. Heselden, J. Chem. SOC., Faraday Trans. I , 1977, 73, 84. Amsterdam, 1983), chap. 2. (PAPER 4/2014)

ISSN:0300-9599    

DOI:10.1039/F19858102155

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I, 1985, 81, 2169-2177 Interfacial Tension Minima in Oil + Water + Surfactant Systems Effects of Salt, Temperature and Alkane in Systems containing Ionic Surfactants BY ROBERT AVEYARD,* BERNARD P. BINKS AND JEREMY MEAD Department of Chemistry, University of Hull, Hull HU6 7RX Received 26th November, 1984 In systems consisting of an n-alkane and a dilute ionic micelle-forming surfactant in aqueous electrolyte it is possible to attain very low oil/water interfacial tensions. These tensions frequently pass through a minimum as the salt concentration, alkane chain length or temperature is varied. We present a thermodynamic treatment of the tension variation and show under what conditions minima can arise. The minimum with respect to salt concentration is seen to result when the effective degree of dissociation of surfactant in micelle and at the oil/water interface are equal (and probably close to zero).At the temperature corresponding to a minimum tension the entropy change on transferring a mole of surfactant (and other associated species) from solution to the oil/water interface is approximately equal to the entropy of formation of micelles containing a mole of surfactant. The occurrence of ultralow interfacial tensions in oil + water systems containing pure ionic surfactant at high dilution is now well recorded, much of the early work having been done by Wade et al.’ Minima in tension with respect to salt concentration and (in the case where the oil is alkane) alkane carbon number are often observed.As with similar systems containing non-ionic surfactant,2 a minimum in tension with respect to temperature is also p~ssible.~ Because of the relevance of ultralow tensions in enhanced oil recovery, much work has also been carried out using commercial mixtures of surfactants (e.g. petroleum s~lphonates).~~ Chan and Shah4 conclude that minimum tension occurs when the equilibrium aqueous phase is at its critical micelle concentration (c.m.c.) and simultaneously the distribution coefficient of surfactant between oil and water is unity; this condition also corresponds to maximum surface excess of surfactant in the monolayer at the oil/water interface. Whatever may be the case for mixed surfactants, theseclaims do not appear to apply to a pure micelle-forming surfactant.We have studied systems containing Aerosol OT (AOT) extensively and summarise our findings later in this paper. A full report will be given el~ewhere.~ Some of the most recent work on dilute pure anionic surfactant systems has involved the use of 8-phenylhexadecane sodium sulphonate (Texas 1) which forms, rather than micelles, dilute dispersions of liquid crystallites at room temperature.s-8 In these systems ultralow tensions have been associated with formation of three-component liquid-crystalline phases at the oil/water interface. It is in the nature of the systems that tensions age considerably and cannot be determined reproducibly. The data are not therefore susceptible to thermodynamic analysis and none has been attempted. Very interesting work has been done on the system sodium dodecyl sulphate+ aqueous NaCl+ toluene + butanol (as ‘co~urfactant’).~~ lo In two-phase regimes Cazabat et u1.l0 believe that ultralow tension is associated with a ‘very thin’ adsorbed surfactant layer.In the three-phase regime, however, tensions between the third 21692170 TENSION MINIMA IN IONIC SURFACTANT SYSTEMS (‘middle’ or ‘surfactant-rich’) phase and one of the outer (oil or aqueous) phases, which are lower than the minimum oil/water tension, are thought to have a quite different origin. Interfacial thicknesses are large and it is appropriate to view these tensions in terms of critical phenomena.lO* l1 Our concern in this paper is with tensions between oil and aqueous phases, whether or not the system is three phase, and we attempt to gain some insight into the origin of tension minima observed with respect to salinity, temperature and alkane carbon number.We report the details of experimental work on alkane + water systems containing AOT and NaCl elsewhere3 but our results can be summarised as follows. (a) At fixed salt concentration and temperature, the oil/water tension falls as surfactant concentration is increased, then levels off at the c.m.c. to the constant value ye. (b) At low salt concentration surfactant resides in the aqueous phase, even above the c.m.c., and at high salt concentration surfactant transfers to the oil phase leaving the aqueous phase close to the c.m.c. with no micelles present. In the latter case the oil phase is a dilute W/O microemulsion. For both high and low salt concentration the tension remains equal to yc once the c.m.c.has been attained in the aqueous phase. The transition between the low and high salt concentration behaviour corresponds to the minimum in yc and it is here that phase-inversion occurs. (c) An analogous situation to that in (b) holds when temperature is varied rather than salt concentration but now surfactant resides largely in the oil phase at low temperature (with the aqueous phase at the c.m.c.) and in water above the c.m.c. at high temperature. We would remark that all the tensions on salt, temperature and alkane ‘scans’ are yc values. Furthermore, the distribution ratio of surfactant between oil and water is seen to be an irrelevance. The tension attains the value yc just as the c.m.c.is attained, so here the distribution ratio K (= surfactant concentration in oil/concentration in water) is zero, and if surfactant then transfers to oil K increases as surfactant is added to the system while y remains constant. If conditions are such that surfactant does not transfer to oil then obviously Kremains zero. In any event the significance of K = 1 depends on the concentration scales used.12 The thermodynamic treatment which follows is based on our findings for micelle- forming surfactants described above and relates to a single interface. It is not relevant therefore to the findings for surfactants forming liquid-crystalline where a liquid-crystalline film intervenes between oil and aqueous phases and two interfaces exist. Ruckensteinl3? l4 has treated similar systems to those of interest here but no simple equations which demonstrate how a tension minimum can occur are given.DISCUSSION We consider a system containing oil (o), water (w) and a 1 : 1 electrolyte (NaCl) having a common cation (Na+) with an anionic surfactant whose anion is designated D-. The variation of the oil/water tension y with surfactant molarity m,- is taken to be as shown in fig. 1, where the constant y attained at the c.m.c. for a given salt concentration and temperature (points A and B) is designated yc. In systems where the aqueous-phase surfactant concentration is initially greater than the c.m.c., the oil/water tension in the equilibrated system is yc. For variations in salt concentration, temperature and, if the oils are n-alkanes, alkane chain length N , yc can pass through a minimum and is frequently ultralow.The object of the present treatment is to obtain expressions for dyc/d In [salt] and dy,/dTand to show under what circumstances these quantities can change sign. We are particularly interested in identifying the important factors associated with tension minima.R. AVEYARD, B. P. BINKS AND J. MEAD 2171 In m, Fig. 1. Schematic representation of y against lnm, for two salt concentrations. We employ the following terms: y is the oil/aqueous-phase interfacial tension and ye is tension at and above c.m.c., m are molar concentrations and a are molar activities, r are the surface excesses relative to T(w) = 0 andf+,f- and f+ are molar activity coefficients of cation, anion and mean ionic activity coefficient, respectively.VARIATION OF ye WITH SALT CONCENTRATION The change in yc between points A and B (fig. 1) brought about by addition of salt dYC d lnm, at constant temperature is given by (1) -- ~ d lnm,a -( a 1nmNa ) D, T + ( & ) N a , T d l n m , , where the signs on the ions have been omitted for clarity. The terms in parentheses on the right-hand side of eqn ( I ) can be obtained from the Gibbs equation, which (2) for constant pressure is -dy = Tidpi+S;dT. i We employ the Gibbs convention for the surface and choose the dividing plane such that the surface excess of water is zero, so that the other r are relative adsorptions and the excess entropy per unit area of surface, S;, is the relative surface e n t r 0 ~ y . l ~ At and below the aqueous phase c.m.c.we may suppose there are negligible amounts of water and surfactant in the oil so we may neglect the term in dp(oi1). Assuming no hydrolysis of the surfactant occurs terms in H+ and OH- may be omitted so that at constant T eqn (2) becomes -dy = RT(T,, d In a,a + T D d In a, + Tcl d In acl). (3) Thus There is experimental evidence that, in systems similar to those of present interest, Tc, is very small and slightly negative,16 as might be expected for a surface containing2172 TENSION MINIMA IN IONIC SURFACTANT SYSTEMS a monolayer of anionic surfactant. For the moment then we will assume that Tc, is effectively zero so that for electroneutrality in the interface TNa = TD. (Later we consider the case where Tcl is not zero.) Thus, eqn (4) may be written wherefTaD is the mean ionic activity coefficient of the surfactant in the presence of the supporting electrolyte, which in cases of interest here is present in large excess.We may reasonably assume that fYaD = fTacl. The second term in parentheses in-eqn (l), obtained from eqn (3) assuming as before that Tcl = 0 and TNa = Tn, is given as In the special case of interest here m, = c.m.c. (in the presence of solubilised alkane).? Also, we may assume that a 1nfYaD/a - In rn, is negligible so that combination of eqn (I), ( 5 ) and (6) yields where changes infYaD have been taken to be equal to changes infYaC1. The salt concentration is equal to mcl, whereas mNa, the total counterion conientration, is equal to the salt concentration+c.m.c. It is widely found that d lnc.m.c./d hm,, is constant.Inspection of eqn (7) shows that for minimum yc ( 3 lnfTacl) + d In c.m.c. = - 1 a In mNa D, T In mNa where both terms on the left-hand side are negative and for sodium chloride in water at 25 "c, - 2(a lnf+ /a In mNa) reaches a maximum value of ca. 0.18 (see fig. 2). Thus, unless d In c.m.c./dln mNa is at least ca. 0.82 no minimum yc appears possible using sodium chloride. For example sodium dodecyl sulphate (SDS) is not expected to give a minimum in yc since with sodium chloride d In c.m.c./d In mNa = - 0.67. In general terms then it appears that whether or not a minimum tension is observed with respect to salt concentration depends crucially on the magnitude of d In c.m.c./d In mNa, but that when a minimum does occur the activity coefficient term is an important factor in determining the salt concentration at which it does so.Basically the shape of the yc against [salt] curve arises from two competing effects on tension when salt is added. At constant m, salt addition lowers y according to eqn (5). On the other hand, rn, is reduced, leading, at constant salt concentration, to an increase in y. The form of the curves predicted by eqn (7) has been obtained as follows. For constant TD and Teqn (7) may be written - yc = RTTD(ln m,, + 2 lnfyaC1 + In c.m.c.) + B'. (9) Values offTaC1 - were taken from ref. (17) and fitted to the cubic in lnm,,: 2 lnf+ - = -6.83 x 10-4(lnm~a)3-0.025(lnm~a)2-0.278 Inn?,,- 1.017. (10) t As discussed elsewhere1* the c.m.c. is the aqueous-phase concentration of surfactant at which surfactant aggregation occurs, in the preferred phase.R. AVEYARD, B.P. BINKS AND J. MEAD 2173 0.09 1 0 0.2 0 .L rn,,/mol dmP3 Fig. 2. Plot of d lnf+/d lnrn,, against rnNa, obtained using data from ref. (17) for NaCl at 25 "C. Furthermore d 1nc.m.c. d lnm,, 1nc.m.c. = a+ ( ) lnmNa where a is a constant, so that eqn (9) becomes - 6.83 x io-yin mNa)3 - o.o25(in mNa)2 d 1nc.m.c. d lnrn,, ) lnmNa]+B (12) where B( = - B'-RTT, a) is the tension at mNa = 1. For the curves shown in fig. 3 a value for rD of 2.28 x mol mP2 is assumed (corresponding to an area per surfactant molecule of ca. 73 A2, which would be appropriate for AOT say). For the purposes of illustration TD is taken as constant but in practice it is expected to decrease at low salt concentrations (below that for which minimum yc is predicted, however).We have arbitrarily set B = 1 mN m-l and curves generated by eqn (12) for three values of d In c.m.c./d In mNa are shown. Curve (a) for d In c.m.c./d In mNa = - 0.7 (similar to that for SDS for example) has no minimum ; a shallow minimum is observed for a value of -0.80. For d lnc.m.c./d lnm,, = -0.85, however, a very sharp minimum is predicted at low salt concentration (mNa z 0.06 mol dmP3) and the curve is very similar to those reported in the literature. Note that the minimum yc for non-ionic surfactants occurs at much higher salt concentrations, as discussed previously. la Further insight into the significance of a minimum in yc can be obtained by noting that Hall19 has shown that - d 1nc.m.c.- a lnf, - (l-%l)+(2-%l)---- d lnm,, a lnm,, where oc, is the effective micellar degree of dissociation defined as twice the negative adsorption of coions and surfactant monomer per micellar surfactant ion. We can similarly define a degree of dissociation ap of a monolayer at a plane surface as21 74 TENSION MINIMA IN IONIC SURFACTANT SYSTEMS 1 0 h I E z E - 1 \ ?- 00 W c - 2 - 3 0 0.25 0 *50 0.75 1.00 ma-/mol dm-j Fig. 3. Curves predicted by eqn (14). TD is taken as 2.28 x mol m-2 (corresponding to an area per molecule of 73 A2) and B as 1 mN m-l. Values of -d lnc.m.c./d lnm,, are (a) 0.70, (b) 0.80 and (c) 0.85. -2rcl/rD (noting that the surface excess of surfactant ions not in the adsorbed monolayer will be negligible in swamping electrolyte).It can be appreciated that by setting Tcl = 0 (the approximation used earlier) we are assuming that the monolayer is completely associated. In the context of the approximation this was justified but as mentioned Tcl is often slightly negative16 and so slight dissociation of the monolayer occurs. If we do not neglect Tcl it is easily shown (noting that m,/mcl 4 1) that in place of eqn (7) we have +%[l+( a lnfzacl ) ]} (14) T D a l n m N a D , T a lnfN,aC1 so that (15) We thus have the interesting result that a minimum in yc is obtained for the salt concentration such that the degree of dissociation of surfactant in the micelle and at the plane surface are equal. On the low-salt-concentration side of the minimum a, > o+,; it is known that micelles are present in the aqueous phase in this salt-concentration regime.On the high-salt-concentration side, however, micelles are absent in the aqueous phase. Since we know that the monolayer at a plane surface is almost completely associated (ap z 0), for micelles to exist in water a, would need to be negative. Physically this implies that the (anionic) micelle would have to take on a positive charge. It can thus be understood that surfactant ‘prefers’ to transfer to the oil phase at higher salt concentrations. It is known that the addition of salt to dilute ionic micellar solutions tends to change the micelle shape in the sequence sphere + rod + disc, i.e. to decrease the average curvature. Presumably at phase inversion and minimum yc the micelle curvature is effectively zero.-~ = R m D { [ i + ( - ) ](a,-,,). d lnmNa a l n m N a D , TR. AVEYARD, B. P. BINKS AND J. MEAD 2175 The balancing factors associated with the micellar surface and the plane monolayer at phase inversion will be seen again in the case where temperature is varied rather than the salt concentration. VARIATION OF yc WITH ALKANE CHAIN LENGTH The salt concentration at which minimum yc is obtained is found to depend on alkane chain length N . It is observed that the higher alkanes give minimum yc at higher salt concentration^.^ For a given salt concentration it is possible for dyc/d lnm,, to be positive for a short-chain alkane and negative for a longer-chain alkane. It follows from eqn (1 7) that a, must be greater for larger chain lengths N . For small N , where dyJd lnm,, is positive, a, < %, implying if % x 0 that a, is negative.Physically this indicates that surfactant transfers to the oil phase leaving the aqueous phase devoid of micelles, as already discussed. In terms of measured quantities the thermodynamic treatment requires that d In c.m.c./d In mNa be numerically larger for systems containing small alkanes. VARIATION OF yc WITH T We finally consider the significance of the way in which yc varies with temperature. For clarity it is convenient to adopt the ‘surface-phase’ model for the interface20 in which total quantities (denoted by superscript s) for the surface appear rather than the excess quantities of the Gibbs convention. We will suppose, as before, that the oil (0) and water (w) are immiscible and we consider for simplicity only the limb of the yc against Tcurve corresponding to the situation where surfactant (D) is present entirely in the aqueous phase.The salt (s) under these conditions will also be entirely in the aqueous phase. The Gibbs equation for this system is -dy = SidT+C. rTdpi (16) i where Ci denotes summation over all the components; the surfactant and salt are considered as the electrically neutral species (D and s). The dp are given by20 and for the pure oil dpo = -SodT where the Si are partial molar entropies of components i in bulk solution and r, and rD are the solute-solvent mole ratios in solution of s and D, respectively.20 For a constant salt concentration the terms in dr, are zero. Furthermore, for very dilute solutions of surfactant in a swamping concentration of electrolyte we may neglect terms in (apu,/ar,) and (apW/arD), so that from eqn (16)-(20) we obtain2176 TENSION MINIMA IN IONIC SURFACTANT SYSTEMS The slope of the yc against Tcurve is given by From eqn (21) (g) = -(s:-z rqsi) r D i Combination of eqn (22)-(24), noting that T , corresponds to the c.m.c., gives (25) -- dye - - (S: -Z rq Si) - RTTS,(d In c.m.c./dT).d T i To a good approximation - RT(d In c.m.c./dT) = ASm, the molar entropy change on forming micelles at the c.m.c., given by19 ASm = Sm-z N i s i (26) i where S, is the entropy of micelles containing one mole of surfactant and Ni is the number of moles of the ith component in the micelles ( N , = 1). Eqn (25), which may shows that the shape of the yc against Tcurve is determined by the relative magnitudes of the entropy of surface formation per unit area (which contains r b moles of surfactant) and the entropy of micelle formation per Tg moles of surfactant.Minimum yc is attained when i.e. when the entropy of transferring a mole of D (with the other associated species) to the plane oil/water interface becomes equal to the molar entropy of micelle formation. Thus, phase inversion and minimum yc are attained, as in the case where the salt concentration is varied, when there is some equivalence between the plane oil/water interface and micelles. We thank British Petroleum for an Extramural Research Award and for the award of a Research Studentship (to B.P.B.). We also thank Prof. D. G. Hall for helpful discussions in connection with eqn (1 5).W. H. Wade, J. C. Morgan, R. S. Schechter, J. K. Jacobson and J. L. Salager, SOC. Pet. Eng. J . , 1978, 242. H. Kunieda and K. Shinoda, Bull. Chem. SOC. Jpn, 1982, 55, 1777. R. Aveyard, B. P. Binks, S. Clark and J. Mead, to be published. K. S. Chan and D. 0. Shah, J . Dispersion Sci. Technol., 1980, 1, 55. M. Dupeyrat, L. Minssieux and A. El Naggar, European Symposium on Enhanced Oil Recovery (Institute of Offshore Engineering, Heriot-Watt University, Edinburgh, 1979), p. 161. E. 1. Franses, J. E. Puig, Y. Talmon, W. G. Miller, L. E. Scriven and H. T. Davis, J. Phys. Chem., 1980,84, 1547. ' E. I . Franses, Y. Talmon, L. E. Scriven, H. T. Davis and W. G. Miller, J . Colloid Interface Sci., 1982, 86,449.R. AVEYARD, B. P. BINKS AND J. MEAD 2177 13 J. E. Puig, E. I. Franses and W. G. Miller, J. Colloid Interface Sci., 1982, 89, 441. A. Pouchelon, J. Meunier, D. Langevin, D. Chateney and A. M. Cazabat, Chem. Phys. Lett., 1980, 76, 277. lo A. M. Cazabat, D. Langevin, J. Meunier and A. Pouchelon, Adv. Colloidlnterface Sci., 1982, 16, 175. l 1 H. T. Davis and L. E. Scriven, Adv. Chem. Phys., 1982,49, 357. l2 R. Aveyard and R. W. Mitchell, Trans. Faraday SOC., 1969, 65, 2645. l 3 E. Ruckenstein and R. Krishnan, J. Colloid Interface Sci., 1980, 76, 201. l 4 E. Ruckenstein, J. Dispersion Sci. Technol., 1981, 2, 1. l5 R. Defay, I. Prigogine, A. Bellemans and D. H. Everett, Surface Tension and Adsorption (Longmans, London, 1966). K. Tajima, Bull. Chem. SOC. Jpn, 1971, 44, 1767. R. A. Robinson and R. H. Stokes, Electrolyte Solutions (Butterworths, Sevenoaks, 1959). R. Aveyard, B. P. Binks, T. A. Lawless and J. Mead, J. Chem. Soc., Faraday Trans. I , 1985,81,2155. Is D. G. Hall, in Aggregation Processes in Solution, ed. E. Wyn-Jones and J. Gormally (Elsevier, Amsterdam, 1983), chap. 2. *O E. A. Guggenheim, Thermodynamics (North Holland, Amsterdam, 5th edn, 1967). (PAPER 4/2015)

ISSN:0300-9599    

DOI:10.1039/F19858102169

出版商:RSC    

年代:1985

数据来源: RSC

摘要:

J . Chem. SOC., Faraday Trans. 1, 1985,81, 2179-2190 Surface Reduction of some Transition-metal Oxides An X-ray Photoelectron Spectroscopic Study of Iron, Cobalt, Nickel and Zinc Oxides BY SIMON J. COCHRAN Department of Chemistry, Monash University, Clayton, Australia 3 168 AND FRANK P. LARKINS* Department of Chemistry, The University of Tasmania, Hobart, Australia 700 1 Received 27th November, 1984 X-Ray photoelection spectroscopy has been used to examine the effect of heating under high-vacuum conditions of some transition-metal oxides. Nickel and cobaltic oxides were found to undergo reduction to Ni metal and COO below 300 "C. This reduction is shown to be caused by reaction with surface carbon-containing contaminants. Removal of this contamination by an in situ oxidation treatment resulted in a well defined surface which was stable to further heating.High-surface-area NiO underwent more extensive reduction than low-surface-area material. In contrast, ZnO was not reduced and only minor changes in the Fe,O, spectra were observed. The results are explained by reference to bulk thermodynamic data. The findings are important particularly for characterisation of oxide surfaces used in catalytic studies. X-Ray photoelectron spectroscopy (X.P.S.) has been widely used to study the surface composition of the transition-metal 0 ~ i d e s . l ~ ~ It is well established that many of the physical properties of the transition-metal oxides depend on their preparation. Yet in the studies of adsorption which have been reported, the catalyst surface has often not been well characterized.In X.P.S., heating in situ is a commonly used method for cleaning sample surfaces. Furthermore, differences in the thermal properties of a substrate may influence its chemisorption properties. Before carrying out any study of adsorption on these catalytically active materials it is important to have a well characterized surface. Therefore, a significant part of the present work has been to characterize the oxide surfaces in order to explain differences in physical properties. Particular emphasis has been placed on nickel oxide, which has been the most widely studied material. A range of samples, both from commercial sources and from preparations in the laboratory, have been studied. This work is intended to provide a foundation for detailed studies of chemisorbed species on various transition-metal-oxide-based catalysts.It has been found that nickel oxide, when prepared by the vacuum decomposition of the hydroxide, degrades to metallic nickel if heated above 250 "C in U ~ C U O . ~ ~ Nickel oxide prepared by vacuum decomposition of the hydroxide forms a catalytically active, high-surface-area material, which readily changes colour from green to black as it takes up excess oxygen on exposure to air. The temperature range over which the active oxide can be produced is quite small (200-220 "C). In contrast, samples prepared by thermal decomposition of nickel salts at high 217921 80 X.P.S. STUDY OF Fe, co, Ni AND Zn OXIDES temperatures in air or oxygen for extended periods show relatively enhanced under ultrahigh-vacuum (u.h.v.) conditions.Furthermore, oxide samples formed by in situ oxidation of metallic nickel are also apparently stable when heated in vacuum, although Schon and Lundin9 reported reduction of surface oxide layers of a nickel sample heated to 300 "C. The oxide prepared by heating in air is of relatively low activity and surface area and remains green on exposure to air. However, Hirokawa et aZ.,lo using x.P.s., have found that NiO samples prepared by the latter method may also decompose, although Dianis and Lester" and Roberts and Smart8 did not report such decomposition when using similar conditions. Nickel oxide is an antiferromagnetic insulator in pure form becoming paramagnetic above the Nee1 temperature (525 K).12 Previous studies have shown that on the surface at least the oxide is rarely if ever stoichiometric, having excess oxygen and behaving as a p-type semiconductor.Co,O, is a dual-valency intrinsic p-type semiconductor and has been studied several times using X.p.s.13-16 It has been found to be one of the most efficient oxide catalysts for the total oxidation of hydrocarbons. It has a spinel structure with C O ~ ~ ' ions occupying the octahedral and C d l the tetrahedral sites.17 Previous workers have differed in the interpretation of .heir X.P.S. results, although there is now general agreement on the origins of the main features. Zinc oxide was chosen to contrast with the preceding oxides. ZnO is an n-type semiconductor with a hexagonal structure. Although it has received some attention in U.P.S.studies, only a small amount of X.P.S. work has been Because of the full 3dsubshel1, its spectrum is much simpler than that of the true transition metals, being characterized by sharp peaks and the absence of shake-up satellites. Iron also forms several oxide phases, which have been studied a number of times by X . P . S . ~ ~ - ~ ~ In this study only Fe20,, an n-type semiconductor, was investigated. EXPERIMENTAL Spectra were recorded on an AEI (Kratos) ESlOO spectrometer, which has been modified by the addition of an u.h.v. sample-preparation chamber and an improved vacuum system, using A1 Ka,, radiation (1486.6 eV) from an X-ray gun operating at 15 kV and 15 mA. Base pressure was better than 2 x Torr (1 Torr z 133.3 Pa) in the analyser and 1 x 10-0 Torr in the sample-preparation chamber after baking the system at 150 "C for 2 days.Normal working pressure in the sample preparation chamber as ca. 2 x Connected to the sample-preparation chamber were a Leybold-Hereaus 4200 quadrupole mass spectrometer (q.m.s.) coupled to a PP20 peak processor for monitoring residual and evolved gases, a Vacuum Generator VG5 argon-ion etching gun, a Barocel capacitance manometer, a metal evaporation facility and a retractable cylindrical quartz oven with gas inlet, capable of heating the sample to 400 "C. Commercial samples of nickel oxide were obtained from Univar (black A.R. grade 'Ni203', N, B.E.T. surface area 160 m2 g-l) and Johnson Matthey (green Specpure grade NiO, surface area 5 m2 g-l). In addition, samples of NiO were prepared by the following methods.(a) In situ decomposition at 200 "C of Ni(OH),, which had been prepared by steam distillation of the he~ammine;~ this material was black and had similar properties to the Univar oxide. (b) Decomposition of Ni(NO,), by heating in air at 600 "C for 18 h. (c) Heating a sample of the Univar oxide in oxygen at 1000 "C for 18 h (surface area < 1 m2 g-l). A sample was also supplied by Smart (Univar oxide heated to 1100 "C, surface area 0.13 m2 g-l) as used in his work.s This sample, as well as (b) and (c), above were green and had similar properties to the Specpure oxide. ZnO, Co,O, and a-Fe,O, were obtained from Johnson Matthey (Specpure grade). ZnO was a fine white powder that became yellowish on pressing or heating in V ~ C U O because of loss of oxygen. Co,O, and Fe,O, were black and red-brown powders, respectively.Total metal ion Torr.S. J. COCHRAN AND F. P. LARKINS 2181 I I 1 865 861 857 853 849 845 535 533 531 529 527 525 binding energy/eV Fig. 1. (A) Ni 2p,,, spectra of high-surface-area NiO: (a) fresh sample, (6) after heating for 1 h at 200 "C, (c) after heating for a further 1 h at 350 "C and ( d ) after heating for 1 h in oxygen at 300 "C and outgassing at 200 "C. (B) 0 Is spectra of high-surface-area NiO recorded immediately after the corresponding spectra in (A). impurities in these samples were < 10 ppm. A wide scan of all samples showed no evidence for any surface contaminant other than carbon to the limit of X.p.s. sensitivity (ca.1 % ). Spectra were calibrated by the vapour deposition of gold (4f7,, binding energy 83.98 eV) on the sample surface.2s Care was taken to use a minimum amount of gold. The C 1s line was not suitable because of its very low intensity and poor definition on samples after the oxidation pretreatment. The 0 1s binding energy of each oxide was found to be constant after correcting for charging and was used as a secondary standard for most experiments. Spectral intensities were corrected using theoretical cross-section~~~ and an experimental analyser sensitivity function.2s Unless otherwise stated, peak areas quoted in this work are those obtained by deconvolution using Gaussian lineshapes after suitable smoothing, background subtraction and satellite removal. RESULTS NICKEL OXIDES The Ni 2p3,, and 0 1s spectra of the high-surface-area Univar nickel oxide are shown in fig.1(A) and (B), respectively. The main spectral features are consistent with previously published Spectrum (b), obtained by heating the fresh sample [spectrum (a)] at 200 "C for 1 h is characterized by lower intensity in the regions of 856.1 eV [fig. 1 (A)] and 531.4 eV [fig. 1 (B)]. These observations correspond to the reduction of Ni3+ ions and the loss of adsorbed water and 0-, respectively. The surface has therefore become more stoichiometric. After heating at 300 "C for 1 h, an intense2182 X.P.S. STUDY OF Fe, Co, Ni AND Zn OXIDES 865 861 857 853 849 845 535 533 531 529 527 525 binding energy/eV Fig. 2. (A) Ni 2p3,2 spectra of low-surface-area NiO: (a) fresh sample, (b) after heating for 1 h at 200 "C and (c) after heating for a further 1 h at 350 "C.(B) 0 1s spectra of low-surface-area NiO recorded immediately after the corresponding spectra in (A). peak at 852.6 eV corresponding to metallic nickel appears as shown in fig. 1 (A)(c). There is also a further loss in intensity of the surface oxygen peak at 531.4 eV and an overall increase in the Ni to 0 intensity ratio. However, after heating the material in oxygen at 300 "C and outgassing, the surface of the material has been reoxidized as shown by spectrum (d). Furthermore, although spectrum ( d ) is similar to spectrum (b) obtained before reduction and subsequent reoxidation, the material is now stable to further heat treatment under vacuum to > 400 "C. A similar result occurred for all the low-surface-area samples.The findings are typified by that of the Specpure material, shown in fig. 2. In this case, however, the extent of reduction is much lower than for the higher-surface-area oxide after identical thermal treatment. All nickel oxide samples examined, with the exception of one prepared by in situ decomposition of the nitrate, showed some degree of reduction after heating at 300 "C under vacuum, the extent depending on the preparative method. Fig. 3 shows the effect of heating on the C 1s spectra of the Specpure oxide. Fig. 3(b) shows that on heating to 200 "C there is some loss of the C 1s intensity relative to the evacuated fresh sample shown in fig. 3(a). The decomposition to the metal is accompanied by a dramatic decrease in the C 1s intensity, resulting in the spectrum of fig.3(c): Heating the sample in 100 Torr of oxygen for 1 h results in the essentially complete removal of this impurity, as shown in fig. 3(d). Similar results were obtained for all the nickel oxide samples investigated. Fig. 4 shows the thermal desorption spectra of the high-surface-area oxide, indicating that the loss of mass is primarily due to removal of surface and lattice oxygen as water and carbon oxides. Very little oxygen gas is desorbed, indicating directS. J. COCHRAN AND F. P. LARKINS 2183 295 291 287 203 279 275 binding energy/eV Fig. 3. C 1s spectra of low-surface-area NiO: (a) fresh sample, (b) after "eating for 1 h at 00 "C, (c) after heating for a further 1 h at 350 "C and ( d ) after heating for 2 h in oxygen at 300 "C and outgassing at 200 "C.0 100 200 T/"C 300 Fig. 4. Thermal-desorption spectrum of contaminants from the surface of a high-surface-area NiO sample. Heating rate: 5 "C min-l. (-) H,O, (-.--) CO,, (---) CO and (...-) 0,.2184 X.P.S. STUDY OF Fe, Co, Ni AND Zn OXIDES 790 786 782 778 774 770 535 533 531 529 527 525 Fig. 5. Fig. 6. Fig. 5. CO 2p,,, specpure Co,O,: (a) fresh sample, (6) after heating for 1 h at 200 "C, (c) after heating for a further 1 h at 350°C and (d) after heating for 1 h in oxygen at 300°C and outgassing at 200 "C. Fig. 6. 0 1s spectra of Specpure Co,O, recorded immediately after the corresponding spectra in fig. 5 . binding energy/eV binding energy/eV decomposition is not significant. A similar result was obtained for the temperature- programmed-desorption spectrum from the low-surface-area oxide.Negligible 0, was again evolved, but significant quantities of CO and CO, in addition to H,O were produced in all cases. Thermogravimetric analysis of the high- and low-surface-area oxides showed that 21% of the initial mass of the high-surface-area oxide was lost on heating in vacuu at 200 "C, whereas only 5% was lost in the case of the low-surface-area oxide. COBALT OXIDE The effect of heating Specpure Co,O, on the Co 2p3,, region of the spectrum is shown in fig. 5. As for the nickel oxide system the Co 2p peaks are broadened by multiplet splitting effects and the CoII and CoI" peaks are not resolved. An asymmetric peak with a maximum at 779.5 eV was obtained for the fresh sample in fig.5(a). On heating to 200 "C, the appearance of a greater intensity on the high-binding-energy side of the main peak in the spectrum shown in fig. 5(b) at 781 eV is indicative of the production of paramagnetic C0V5 Note that the binding energy of CoI1 in this case is higher than that of CoIII, which is the opposite to the result normally expected. This is believed to be due to the fact that the CoI1 ions occupy tetrahedral sites inS . J . COCHRAN AND F. P. LARKINS 2185 a high-spin configuration, with a larger number of close neighbours than Co"' in the octahedral sites, which is low spin and diamagnetic. On heating at 350 "C, a pronouned CoII satellite appears at 786 eV, as can be seen in fig. 5(c). This indicates extensive reduction of the sample surface to COO. No evidence was obtained for metallic Co under these conditions.Following heating in oxygen, the surface is reoxidized, as shown in fig. 5(d), and as in the case of NiO remains stable to further heating in uacuo. Deconvolution of fig. 5 ( d ) into two peaks indicates binding energies of 779.4 and 780.3 eV for ColI1 and Co", respectively. The surface C O I ~ ~ / C O ~ ~ ratio is ca. 4.2, compared with an expected value of 2.0. The corresponding 0 1s spectra are shown in fig.6. As in the case of nickel oxide there is a peak due to lattice oxygen at 529.6 eV in all the spectra and a smaller peak at 53 1 eV due to adsorbed H,O and OH-, with intensity strongly dependent on sample history. The effect of heating on Co,O, [shown in fig.6(b) and (c)] is also similar to that of NiO. In both cases there is a decrease in the intensity of the non-stoichiometric oxygen peak as the preheating temperature is raised. Fig. 6(d) shows that heating in oxygen followed by outgassing in uacuo does not remove all the excess surface oxygen. Fig. 7 shows the C Is region of the spectra recorded immediately after the Co and 0 Is spectra. After heating at 350 "C the dramatic reduction of the C I s intensity in fig. 7(c) can be observed relative to fig. 7(b). Little improvement was obtained after heating in oxygen, as fig. 7(d) shows. The reduction of Co,O, to COO appears even more facile than the reduction of NiO to Ni. This is supported by thermodynamic evidence, which will be discussed later. Thermal desorption from Co,O, also shows the loss of carbon-containing species in a similar manner to nickel oxide.ZINC OXIDE No changes were observed in the Zn 2p or L,M,,,M, Auger spectra on heating or oxidation of ZnO. The spectra are not presented here as they do not differ from earlier work.18 Although, being an n-type material, one might have expected the presence of ZnO under the present experimental conditions, this species was not detected by X.P.S. The binding energy shift of the Zn 2p,/, lines between the oxide and metal has been estimated to be < 0.5 eV;18 however, the L, M4, , M4, , Auger shift is ca. 4.3 eV and would have been readily observed if metallic zinc was present. By comparison with the nickel and cobalt oxides, surface carbon contamination on zinc oxide was much less readily removed by heating.Several oxidation treatments at 300 "C were sometimes necessary to achieve its complete removal. The finding is consistent with a mechanism whereby the lattice oxygen ions are not significantly involved in the oxidation of the surface carbon species. IRON(III) OXIDE The effect of heating at 350 "C in uacuo for 4 h on the Fe 2p3/, spectrum is shown in fig. 8. Only slight evidence for the production of FeI1 was found. The Fe 2p3/, peak broadened slightly relative to the fresh sample and there was an increase in satellite intensity at ca. 7 18 eV. No metallic iron was observed even after extensive heating. The thermal desorption spectrum revealed significant quantities of CO and CO, being evolved in addition to some H,O.This observation, coupled with X.P.S. evidence for the removal of surface carbon on heating, is consistent with the reduction of Fe,O, to Fe,04 under the experimental conditions. The absence of definite X.P.S. evidence for Ferl may be explained either by the closeness of the FeI1 and FelI1 binding energies or by the diffusion of FeI1 species to the bulk. Suitable data for comparison with related inorganic compounds are not available.2186 X.P.S. STUDY OF Fe, Co, Ni AND Zn OXIDES 200 counts A 1 I 1 I 296 292 288 204 280 276 740 732 724 716 708 700 binding energy/eV binding energy/eV Fig. 7. Fig. 8. Fig. 7. C 1s spectra of Specpure Co,O, recorded after the corresponding spectra in fig. 5. Fig. 8. Fe 2p3,2 spectra of Specpure Fe,O,: (a) fresh sample after evacuation and (6) after heating at 350 "C.DISCUSSION To interpret the observations the standard Gibbs free energy values for the oxides were examined using the CSIRO Thermodata base.30 The thermodynamic data are based upon the JNAF tables and literature sources. Free energy values AGe( T ) have been calculated for a series of relevant reactions at temperatures up to 400 "C. These calculations are for what must be considered model reactions indicative of the processes occurring at the oxide surface. Results are presented in table 1. The AGe values can be for guidance only since under reaction conditions in the ultrahigh-vacuum chamber partial pressures of gaseous reactants and products are very low. Furthermore, no insight concerning the kinetics of the various processes is provided by consideration of thermodynamic data.It is clear from table 1 (A) that all the oxides studied should be thermally stable against decomposition at temperatures frequently used in catalytic processes. To explain the decomposition of high-surface-area nickel oxide, it was previously proposed that a highly strained environment and a relatively large proportion of surface nickel atoms were the causes of the red~ction.~ Indeed, it is known that bulk thermodynamics is often unsuitable for explaining the behaviour of small particle~.~~ However, the generality of the effect has been demonstrated here for a range of oxides with different histories. The lack of evolved oxygen and the evidence that reoxidized carbon-free samples do not decompose supports the alternativeS.J. COCHRAN AND F. P. LARKINS 2187 Table 1. Standard free energies for various oxide reactions (in kJ mol-l) as a function of temperature temperature/'C reaction 150 200 250 300 3 50 400 (A) decomposition NiO(s) + Ni(s)+$O,(g) 200.8 196.2 191.7 187.3 C0304(s) + ~COO(S) +$O,(g) 133.5 125.4 117.1 108.7 3Fe203(s) + 2Fe304(s) + iodg) 179.5 173.0 166.5 160.0 ZnO(s) + Zn(s)+$O,(s) 305.8 300.9 295.9 291.0 NiO(s) + C(s) + Ni(s) + CO(g) 52.4 43.3 34.2 25.3 Co304(s) + 4C(s) + 3Co(s) + 4CO(g) 153.0 115.5 78.1 40.7 Fe,O,(s) + 3C(s) + 2Fe(s) + 3CO(g) 264.4 237.4 210.5 183.8 ZnO(s)+C(s) + Zn(s)+CO(g) 157.4 147.9 138.4 129.0 (B) carbon reduction, carbon monoxide product CO~O~(S)+C(S) -P ~COO(S)+CO(~) - 14.9 -27.6 -40.4 -53.4 3Fe,03(s) + C(s) + 2Fe304(s) + CO(g) 3 1 .O 20.0 9.0 - 2.0 (C) carbon reduction, carbon dioxide product 2NiO(s) + C(s) + 2Ni(s) + CO,(g) 6.8 -2.4 - 11.6 -20.6 ~CO~O~(S)+C(S) + ~COO(S)+CO,(~) - 127.7 - 144.1 - 160.8 - 177.7 CO~O~(S)+~C(S) + 3Co(~)+2CO,(g) -42.9 -62.5 -82.0 - 101.5 6Fe,O,(s)+C(s) +4Fe304(s)+C02(g) -35.8 -48.9 -61.9 -75.1 2Fe,O,(s) + 3C(s) -, 4Fe(s) + 3CO,(g) 1 17.5 103.9 90.4 77.1 2Zn0 + C(s) + 2Zn(s) + CO,(g) 2 17.0 206.9 196.9 186.9 182.9 100.2 153.4 286.1 16.3 -66.3 3.3 - 13.2 157.2 119.6 - 29.5 - 194.7 - 121.0 - 88.4 63.9 177.0 178.4 91.7 146.7 281.2 7.4 - 79.4 - 33.9 -24.3 130.7 1 10.2 - 38.6 -21 1.9 - 140.4 - 101.9 50.9 167.1 explanation, namely that reduction is caused by reaction with surface carbon contaminants. Although it has previously been suggested that nickel oxide may be reduced by reaction with carbon-containing c~ntamination,~~ this hypothesis has been The standard-free-energy values for the reduction of the oxides using carbon as reductant yielding carbon monoxide or carbon dioxide as the product are shown in table l(B) and (C).The data show that nickel oxide is spontaneously reduced by carbon at temperatures > 200 "C with carbon dioxide rather than carbon monoxide as the major gaseous carbon oxide product. This is consistent with the observations reported in fig. 4. While AGe is positive for the production of carbon monoxide, in view of the low partial pressure of CO in the chamber (background estimate < 6 x lo-* Pa), AG may also be negative. Hence, under u.h.v. conditions, CO as well as CO, may be produced.The AGe value for the reduction of c0304 to COO by carbon predicts that the process is spontaneous for the production of both CO and CO,. Further reduction to cobalt metal is predicted to be favoured only with the production of CO,. No evidence for cobalt metal was found in the experimental observations. It is probable that insufficient carbon-containing impurity was available to induce the two-stage reduction. The data in table 1 are consistent with the reduction by carbon of the oxide Fe203 to Fe304, principally to yield carbon dioxide as the major gaseous product, but some carbon monoxide is also expected. While the FeIL and FelI1 2p photoelectron peaks could not be readily distinguished in the spectrum, the gases evolved were consistent with the thermodynamic predictions.No evidence was obtained for the further reduction to metallic iron. This is consistent with the thermodynamic data. Zinc oxide is not predicted to be reduced under the conditions used and this appears to be the case experimentally. and the evidence previously has been largely circumstantial.2188 X.P.S. STUDY OF Fe, Co, Ni AND Zn OXIDES A hypothesis that the backgroundpressure of CO in the spectrometer is responsible for the reduction cannot account for the loss of carbon from the sample surface and the increase in the partial pressure of CO, which is observed as the sample is heated. Furthermore, the partial pressure of CO in the preparation chamber is estimated to be < 6 x lop8 Pa. Thermodynamic data calculated but not presented here supports the view that if sufficient carbon monoxide was present it would reduce NiO, Co,O, and Fe,O,, but not ZnO.The production of significant quantities of H20 in the thermal-desorption spectra of the oxides is consistent with hydrocarbon-type contaminants having a role in the reduction process. For illustrative purposes the AGO values for the reaction of methane with nickel oxide 4NiO(s) + CH,(g) -P 4Ni(s) + CO,(g)+ 2H,O(g) were considered. AG* values were negative, indicative of a spontaneous reaction for temperatures > 175 "C. Another source of H20 in the spectrum is due to the direct removal of strongly bound water and surface hydroxyl groups. The nature of the contamination is not easy to establish. In earlier work by many researchers using poorer vacuum, backstreaming oil from the diffusing pump was generally believed to be the major source of contamination.However, this is not the case in the present study. For the high-surface-area oxide prepared at low temperatures, it is undoubtedly introduced during preparation and is impregnated throughout the oxide. In the case of the low-surface-area oxides, however, it would be expected that during high-temperature heating in air, any carbon present would be oxidized and removed from the sample, with immediate reoxidation of any reduced species formed. Likewise, after oxidizing in situ very little carbon is accumulated on the surface. Thus, the major contamination would appear to occur on the sample surface and arise from exposure to the atmosphere during storage and pressing of the samples.Visual examination of the surfaces of the samples indicate contamination of different origin. Whereas high-surface-area nickel oxide evidently underwent a bulk decomposition and turned completely black, reduction of low-surface-area samples was confined to the surface, which acquired a greyish appearance. This reinforces the view that carbon is distributed throughout the bulk of the high-surface-area oxide. It has been previously established with other catalyst systems that the lattice oxygen of a reducible metal oxide may act as a useful oxidizing agent for hydrocarbons. The present work indicates that non-stoichiometric high-surface-area nickel oxides are more readily reduced and more likely to be an active catalyst than the relatively stoichiometric oxides.Further elucidation of the role of lattice oxygen in oxidation and reduction reactions may be undertaken with isotopic studies. Because of the involvement of lattice oxygen in surface reactions, any thermodynamic treatment of such reactions must therefore consider the thermodynamics of the substrate. While several of the more recent infrared studies of adsorbates on oxide surfaces have involved pretreatment by heating in oxygen, this practice has rarely been performed in X.P.S. studies. This procedure should be adopted in order to simplify the interpretation of results and to improve reproducibility. This work demonstrates that while heating alone is at least partially successful in cleaning the surface, a significant change in the chemical nature of the material may occur.Argon-ion bombardment is an alternative method of 'cleaning' the surfaces of oxide samples. Following this treatment, reduction is also frequently observed. In the study of reduction of NiO by argon-ion bombardment by Ostyn and Carter,34 reduction was interpreted as being due to differential oxygen sputtering. However, it is noted that ( I ) the temperature at which reduction occurred may have been as high as 200-300 "C,S. J. COCHRAN AND F. P. LARKINS 2189 (2) relatively poor vacuum conditions were used and (3) consistent results were not always obtained for different samples. This strongly suggests that carbon may have been involved in the reduction. Kim et al.35 have presented thermodynamic evidence for the reduction of metal oxides by argon ion bombardment.They found that oxides with heats of formation < -250 kJ mol-l could be reduced, while those with heats of formation > ca. -490 kJ mol-l were not reduced. Since the materials with lower heats of formation would also be more easily reduced by carbon, a chemical reduction involving carbon cannot be ruled out. Note3s that the heat of formation of CO, at - 349 kJ mol-l lies between the two values cited above. CONCLUSIONS It has been shown that the surface reduction induced by heating above 200 "C of NiO and Co,O, surfaces is due to the reaction of the oxides with carbon-containing contamination, while ZnO is not reduced, as predicted by bulk thermodynamic data. The reduction of Fe,O, was not unambiguously determined. High-surface-area nickel oxide was more readily reduced than low-surface-area samples. After oxidation treatment of all the oxides a clean, reproducible sample surface could be obtained which was stable to further treatment.Prior to oxidation, however, surface composition and reducibility were variable, depending on sample preparation and treatment. It is therefore recommended that any study involving oxides using surface-sensitive techniques should involve pretreatment by heating in oxygen. Heating in uacuo or argon-ion bombardment was found to be insufficient treatment for sample cleaning and led to modification of the surface. Financial assistance from the Australian Research Grants Scheme for this research is gratefully acknowledged. We also thank Mr M. Wadsley, CSIRO, for providing the thermodynamic data and the Australian Government for a scholarship (to S.J. C). A. Rosencwaig, G. K. Wertheim and H. J. Guggenheim, Phys. Rev. Lett., 1971,27,479. T. Novakov and R. Prins, in Electron Spectroscopy, ed. D. A. Shirley (North Holland, Amsterdam, 1972). L. Fiermans, R. Hoogewijs and J. Vennik, Surf. Sci., 1975, 47, 1. S. J. Teichner and J. A. Morrison, Trans. Faraday Soc., 1955,51, 961. F. P. Larkins and P. J. Fensham, Nature (London), 1967, 215, 1268. J. Deren and J. Stoch, J . Catal., 1970, 18, 249. ' R. B. Fahim and A. I. Abu-Shady, J . Catal., 1970, 17, 10. a M. W. Roberts and R. St. C. Smart, Chem. Phys. Lett., 1980,69, 234. lo K. Hirokawa, F. Honda and M. Oku, J . Electron Spectrosc. Relat. Phenom., 1975, 6, 333. l 2 I.G. Austin, A. J. Springthorpe, B. A. Smith and C. E. Turner, Proc. Phys. SOC. (London), 1967,90, l 3 J. Grimbolt, A. D'Huysser, J. P. Bonnelle and J. P. Beaufils, J. Electron Spectrosc. Relat. Phenom., l4 M. Oku and K. Hirokawa, J . Electron Spectrosc. Relat. Phenom., 1975, 8, 475. l5 T. J. Chuang, C. R. Brundle, and D. W. h c e , Surf. Sci., 1976, 59,413. l 6 J. Haber and L. Ungier, J . Electron. Spectrosc. Relat. Phenom., 1977, 12, 305. G. Schon, and S. T. Lundin, J . Electron Spectrosc. Relat. Phenom., 1972, 1, 105. W. P. Dianis and J. E. Lester, Surf. Sci., 1974, 43, 602. 157. 1975, 6, 71. F. A. Cotton and G. Wilkinson, Adcanced Inorganic Chemistry (Interscience, New York, 4th edn, 1980). G. Schon, J . Electron Spectrosc. Relat. Phenom., 1973, 2, 75. J. Haber, J. Stoch and L. Ungier, J . Electron Spectrosc. Relat. Phenom., 1976, 9, 459. G. C. Allen, M. T. Curtis, A. J. Hooper and P. M. Tucker, J . Chem. Soc., Dalton Trans., 1974, 14, 1525. 2o M. J. Dreiling, Surf. Sci., 1978, 71, 231.2190 X.P.S. STUDY OF Fe, CO, Ni AND Zn OXIDES 22 K. Asami, K. Hashimoto and S. Shimodaira, Corrosion Sci., 1976, 16, 35. 23 C. R. Brundle, T. J. Chang and K. Wandelt, Surf. Sci., 1977, 68, 459. 2q M. Oku and K. Hirokawa, J. Appl. Phys., 1979,50,6303. 25 K. Wandelt, Surf. Sci. Rep., 1982, 2, 1. 26 R. J. Bird and P. Swift, J. Electron Spectrosc. Relat. Phenom., 1980, 21, 227. 27 J. H. Scofield, J. Electron Spectrosc. Relat. Phenom., 1976, 8, 129. 28 A. Lubenfeld, M S c . Thesis (Monash University, 1977). 29 K. S. Kim and N. Winograd, Surf: Sci., 1974, 43, 625. 30 A. G. Turnbull, CSIRO-NPL Thermodata System, Version 3 Rep. MCC 23, 1981. 31 T. L. Hill, Thermodynamics of Small Systems (Benjamin, New York, 1963). 32 M. E. Dry and F. S. Stone, Discuss. Faraday Soc., 1959,243, 192. 33 S. J. Teichner, Discuss. Faraday SOC., 1959, 28, 215. 34 K. M. Ostyn and C. B. Carter, SurJ Sci., 1982, 121, 360. 35 K. S. Kim, W. E. Baitinger, J. W. Amy and N. Winograd, J. Electron Spectrosc. Relat. Phenom., 1974, 36 CRC Handbook of Chemistry and Physics, ed. R. C. Weast (CRC Press, Boca Raton FL, 61 st edn, 1980). 5, 351. (PAPER 4/2025)

ISSN:0300-9599    

DOI:10.1039/F19858102179

出版商:RSC    

年代:1985

数据来源: RSC