摘要:

4369 4377 4387 4397 4407 4417 4427 4439 445 1 4457 447 1 4475 4487 4495 450 1 4509 Con tents A New Form of the High-temperature Isopiestic Technique and its Applica- tion to Mercury-Bismuth, Mercury-Cadmium, Mercury-Gallium, Mercury- Indium and Mercury-Tin Binary Amalgams Z-C. Wang, X-H. Zhang, Y-Z. He and Y-H. Bao The Derivation of Chemical-diffusion Coefficients of Oxygen in UO,,, over the range 180-300 "C. Spectroscopic Procedure and Preliminary Results T. R. Griffiths, H. V. St. Aubyn Hubbard, G. C. Allen and P. A. Tempest Pho tophysics at Solid Surfaces. Evidence of Dimer Formation and Polarization of Monomer and Excimer Fluorescences of Pyrene in the Adsorbed State on Silica-gel Surfaces T. Fujii, E. Shimizu and S. Suzuki Ordering in Monodispersed Polymer Latices induced by a Temperature Gradient K.Furusawa, N. Tobori and S. Hachisu X-Ray Diffraction Study of Molten Eutectic LiF-NaF-KF Mixture K. Igarashi, Y. Okamoto, J. Mochinaga and H. Ohno Viscosity Measurements of Some Tetra butylammonium, Copper( I), Silver( I) and Thallium( 1) Salts in Acetonitrile-Pyridine Mixtures at 15, 25 and 35 "C D. S. Gill and B. Singh The Ethane- 1,2-diol-Water Solvent System. The Dependence of the Dis- sociation Constant of Picric Acid on the Temperature and Composition of the Solvent Mixture Silver(1) Complexation with Tertiary Amines in Toluene M. Soledade Santos, E. F. G. Barbosa and M. Spiro Enhanced Oxygen Evolution through Electrochemical Water Oxidation mediated by Polynuclear Complexes embedded in a Polymer Film G. J. Yao, A. Kira and M. Kaneko Nature of Acid Sites in SAP05 Molecular Sieves.Part 1.-Effects of the Concentration of Incorporated Silicon C. Halik, J. A. Lercher and H. Mayer Hemimicelle Formation of Cationic Surfactants at the Silica Gel-Water Interface T. Gu, Y. Gao and L. He Nuclear Magnetic Resonance Relaxation in Micelles. Deuterium Relaxation at Three Field Strengths of Three Positions on the Alkyl Chain of Sodium Dodecyl Sulphate Studies of the Temperature Dependence of Retention in Supercritical Fluid Chromatography K. D. Bartle, A. A. Clifford, J. P. Kithinji and G. F. Shilstone Hydrogen and Muonium Atom Adducts of Trimethylsilyl Derivatives of Ethyne The Radical Cation of Formaldehyde in a Freon Matrix. An Electron Spin Resonance Study Phase Transition of the Water confined in Porous Glass studied by the Spin- probe Method H.Yoshioka G. C. Franchini, A. Marchetti, L. Tassi and G. Tosi 0. Soderman, G. Carlstrom, U. Olsson and T. C. Wong C. J. Rhodes and M. C. R. Symons C. J. Rhodes and M. C. R. Symons4369 4377 4387 4397 4407 4417 4427 4439 445 1 4457 447 1 4475 4487 4495 450 1 4509 Con tents A New Form of the High-temperature Isopiestic Technique and its Applica- tion to Mercury-Bismuth, Mercury-Cadmium, Mercury-Gallium, Mercury- Indium and Mercury-Tin Binary Amalgams Z-C. Wang, X-H. Zhang, Y-Z. He and Y-H. Bao The Derivation of Chemical-diffusion Coefficients of Oxygen in UO,,, over the range 180-300 "C. Spectroscopic Procedure and Preliminary Results T. R. Griffiths, H. V. St. Aubyn Hubbard, G. C. Allen and P. A. Tempest Pho tophysics at Solid Surfaces.Evidence of Dimer Formation and Polarization of Monomer and Excimer Fluorescences of Pyrene in the Adsorbed State on Silica-gel Surfaces T. Fujii, E. Shimizu and S. Suzuki Ordering in Monodispersed Polymer Latices induced by a Temperature Gradient K. Furusawa, N. Tobori and S. Hachisu X-Ray Diffraction Study of Molten Eutectic LiF-NaF-KF Mixture K. Igarashi, Y. Okamoto, J. Mochinaga and H. Ohno Viscosity Measurements of Some Tetra butylammonium, Copper( I), Silver( I) and Thallium( 1) Salts in Acetonitrile-Pyridine Mixtures at 15, 25 and 35 "C D. S. Gill and B. Singh The Ethane- 1,2-diol-Water Solvent System. The Dependence of the Dis- sociation Constant of Picric Acid on the Temperature and Composition of the Solvent Mixture Silver(1) Complexation with Tertiary Amines in Toluene M.Soledade Santos, E. F. G. Barbosa and M. Spiro Enhanced Oxygen Evolution through Electrochemical Water Oxidation mediated by Polynuclear Complexes embedded in a Polymer Film G. J. Yao, A. Kira and M. Kaneko Nature of Acid Sites in SAP05 Molecular Sieves. Part 1.-Effects of the Concentration of Incorporated Silicon C. Halik, J. A. Lercher and H. Mayer Hemimicelle Formation of Cationic Surfactants at the Silica Gel-Water Interface T. Gu, Y. Gao and L. He Nuclear Magnetic Resonance Relaxation in Micelles. Deuterium Relaxation at Three Field Strengths of Three Positions on the Alkyl Chain of Sodium Dodecyl Sulphate Studies of the Temperature Dependence of Retention in Supercritical Fluid Chromatography K. D. Bartle, A. A. Clifford, J. P. Kithinji and G. F. Shilstone Hydrogen and Muonium Atom Adducts of Trimethylsilyl Derivatives of Ethyne The Radical Cation of Formaldehyde in a Freon Matrix. An Electron Spin Resonance Study Phase Transition of the Water confined in Porous Glass studied by the Spin- probe Method H. Yoshioka G. C. Franchini, A. Marchetti, L. Tassi and G. Tosi 0. Soderman, G. Carlstrom, U. Olsson and T. C. Wong C. J. Rhodes and M. C. R. Symons C. J. Rhodes and M. C. R. Symons

ISSN:0300-9599    

DOI:10.1039/F198884FX001

出版商:RSC    

年代:1988

数据来源: RSC

摘要:

NOMENCLATURE AND SYMBOLISM Units and Symbols. The Symbols Committee of The Royal Society, of which The Royal Society of Chemistry is a participating member, has produced a set of recommendations in a pamphlet 'Quantities, Units, and Symbols' (1 975) (copies of this pamphlet and further details can be obtained from the Manager, Journals, The Royal Society of Chemistry, Burlington House, London W1V OBN). These recommendations are applied by The Royal Society of Cemistry in all its publications. Their basis is the 'Systeme International d'Unit6s' (9). A more detailed treatment of units and symbols with specific application to chemistry is given in the IUPAC Manual of Symbols and Terminology for Physicochemical Quantities and Units (Pergamon, Oxford, 1979). Nomenclature. For many years the Society has actively encouraged the use of standard IUPAC nomenclature and symbolism in its publications as an aid to the accurate and unambiguous communication of chemical information between authors and readers.In order to encourage authors to use IUPAC nomenclature rules when drafting papers, attention is drawn to the following publications in which both the rules themselves and guidance on their use are given: Nomenclature of Organic Chemistry, Sections A, B, C, D, E, F, and H (Pergamon, Oxford, 1979 edn). Nomenclature of Inorganic Chemistry (Butterworths, London, 1971 , now published by Pergamon). Biochemical Nomenclature and Related Documents (The Biochemical Society, London, 1978). Compendium of Chemical Terminology: IUPAC Recommendations (Blackwells, Oxford, 1987).A complete listing of all IUPAC nomenclature publications appears in the January issues of J. Chem. SOC., Faraday Transactions. It is recommended that where there are no IUPAC rules for the naming of particular compounds or authors find difficulty in applying the existing rules, they should seek the advice of the Society's editorial staff. (xiv)NOMENCLATURE AND SYMBOLISM Units and Symbols. The Symbols Committee of The Royal Society, of which The Royal Society of Chemistry is a participating member, has produced a set of recommendations in a pamphlet 'Quantities, Units, and Symbols' (1 975) (copies of this pamphlet and further details can be obtained from the Manager, Journals, The Royal Society of Chemistry, Burlington House, London W1V OBN). These recommendations are applied by The Royal Society of Cemistry in all its publications.Their basis is the 'Systeme International d'Unit6s' (9). A more detailed treatment of units and symbols with specific application to chemistry is given in the IUPAC Manual of Symbols and Terminology for Physicochemical Quantities and Units (Pergamon, Oxford, 1979). Nomenclature. For many years the Society has actively encouraged the use of standard IUPAC nomenclature and symbolism in its publications as an aid to the accurate and unambiguous communication of chemical information between authors and readers. In order to encourage authors to use IUPAC nomenclature rules when drafting papers, attention is drawn to the following publications in which both the rules themselves and guidance on their use are given: Nomenclature of Organic Chemistry, Sections A, B, C, D, E, F, and H (Pergamon, Oxford, 1979 edn). Nomenclature of Inorganic Chemistry (Butterworths, London, 1971 , now published by Pergamon). Biochemical Nomenclature and Related Documents (The Biochemical Society, London, 1978). Compendium of Chemical Terminology: IUPAC Recommendations (Blackwells, Oxford, 1987). A complete listing of all IUPAC nomenclature publications appears in the January issues of J. Chem. SOC., Faraday Transactions. It is recommended that where there are no IUPAC rules for the naming of particular compounds or authors find difficulty in applying the existing rules, they should seek the advice of the Society's editorial staff. (xiv)

ISSN:0300-9599    

DOI:10.1039/F198884BX003

出版商:RSC    

年代:1988

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1988, 84(1), 9-18 The Dissolution of Magnetite by Nitrilotriacetatoferrate(I1) Margarita del Valle Hidalgo and Nkstor E. Katz Catedra de Fisicoquimica III, Facultad de Bioquimica, Quimica y Farmacia, Universidad Nacional de Tucuman, Ayacucho 491, 4000 Tucuman, Argentina Albert0 J. G. Maroto and Miguel A. Blesa* Departamento Quimica de Reactores, Comisidn Nacional de Energia A tdmica, Avenida del Libertador 8250, 1429 Buenos Aires, Argentina The dissolution of magnetite particles in solutions containing nitrilo- triacetatoferrate(I1) has been studied as a function of total nitrilotriacetic acid (NTA) and iron(1r) concentrations, pH and temperature. Experimental results are interpreted in terms of adsorption equilibria involving the free ligand and its metal complexes, and inner- and outer-sphere interfacial electron transfer from the adsorbed electroactive FeI' species to surface -Fell1 sites, either free or complexed by NTA.Ion transfer-controlled dissolution was ruled out by the experimental evidence. Fey- ions Ty = N(CH,CO,)i-] have been identified as the electroactive species, electron transfer being an outer-sphere process for the present experimental conditions : the precursor complex can be represented by -Fe"'-NTA. . . Fey-. Dissolution rate deviates from first order in [Fey-]; the order decreases with increasing [Fey-] as a consequence of the relatively high affinity of the complex for -Fe"'-NTA surface sites. Reaction order on proton concentration is 0.67, reflecting the requirement of H+ ions adjacent to the site where the reductant is adsorbed.The apparent activation energy is 73 kJ mol-l, which is a composite of equilibrium and kinetic parameters. In spite of their technological importance, metal oxide dissolution processes are by no means well understood.' Different preparations of a given metal oxide respond in various ways to aqueous solvents; surface details and bulk defects play a vital role in determining their reactivity and it is not surprising that there is no adequate (even qualitative) description of the oxide parameters that influence solubility kinetics. Dissolution rates are of course also sensitive to solution parameters; even though in this case the basic ideas can in principle be tackled more easily, there is a noticeable scarcity of detailed studies.Typical solvents for iron oxides are aqueous acid solutions of chelating agents, such as polycarboxylic acids and amino acids.2 The effectiveness of such formulations is limited, and there is little evidence on the mechanism of iron leaching into these solutions. Chemisorption of the ligands is involved,l-1° but it is not clear whether the complexed surface -Fe"I ions are intermediates in the dissolution process, or 'dead-end' products that must return to a more reactive uncomplexed state prior to dissolution.lY8 It is even possible that phase transfer of FeIII is negligibly slow in acidic solutions, and that a redox mechanism is inv01ved.~~~ Reductants greatly enhance the rate of dissolution of iron(1n) oxides ; this phenomenon has been known for yearsl1-l4 and involves formation of the more labile iron@) ions by electron transfer across the interface,15-19 followed by iron(I1) ion transfer ; this effect may even be observed in the dissolution in mineral acids.l' Of particular interest is the reductive dissolution by metal complexes, such as V(pic), (pic = picolinate anion)15- l8 or Fe" in solutions containing polycarboxylates.In these latter systems, several concurrent chemical reactions take place simultaneously, and it is not easy to identify 910 Dissolution of Magnetite by Nit rilo tr iace ta toferra te( 11) unambiguously the reasons for the dependence of the dissolution rate on the concentrations of ligand, metal ion and hydrogen ion. The present paper reports the results of a kinetic study of the dissolution of magnetite in solutions containing nitrilotriacetic acid (NTA)* and iron(r1) salts.Together with the results obtained in oxalic acid16* 2o and EDTA (ethylenediaminetetra-acetate)'. 21 media, this information is used to propose a general mechanism for the dissolution of magnetite by iron(I1) ions in the presence of polycarboxylates, in which the relative importance of inner- and outer-sphere electron-transfer processes is taken into account. Experimental Magnetite (Fe,O,) was prepared as described in a previous paper22 by oxidising an iron(I1) salt in an alkaline medium in the presence of hydrazine. It was characterised by chemical analysis, X-ray diffraction, Mossbauer spectroscopy, scanning electron microscopy and surface area measurements.It was composed of cubo-octahedral particles of 0.3 pm average diameter. The oxide was crystalline and stoichiometric, and had a specific surface area of 4.7 m2 g-l. Kinetic experiments were performed in a magnetically stirred and tightly stoppered cylindrical beaker provided with a thermostatted water jacket. Solutions of nitrilo- triacetic acid of suitable pH and concentration, containing the desired amount of Fe'' in the form of (NH,),Fe(S0,),.6H20, were carefully deaerated by bubbling with purified nitrogen. Dissolution was started by pouring magnetite into the solution. In order to avoid oxidation of iron(II), periodic sampling was not employed; instead, the kinetics were followed by measuring the acid consumption necessary to keep the pH constant as a function of time.The pH-stat (Mettler DK-10) was calibrated with three standard buffer solutions. All reagents were analytical grade. Water was twice distilled in a quartz apparatus. Some experiments were carped out as far as complete dissolution. In these cases, the linearity of plots of 1 - (1 - f ) s against time, f being the fraction of the solid remaining at time t, showed a dependence of rate on residual surface area. In general, however, data were analysed in terms of initial rates; accordingly, unless stated otherwise, square brackets represent initial concentrations. Acid consumption data were transformed into rates R, expressed as the number of moles of iron dissolved per unit area and time, by means of eqn (1) R = (3N/nA)(AV/At), (1) and the stoichiometry given by eqn (2) Fe,O, + 2(Fe"HZYy)(3y-z-2)- + qH,Y(3-p)- + nH+ In eqn ( l ) , N is the molarity and AV the volume of monobasic titrant; A is the total available surface area (A/m2 = 4.7 x w, w being the mass in g of magnetite) and n is the number of protons as defined in eqn (2).Eqn (2) was solved for n through mass balance and charge equations, using stability constants23 to characterise the average a, b, p, x and y values. Solutions of eqn (2) were checked in selected experiments by measuring n through the simultaneous determination of dissolved iron and consumed acid. * In this paper, the following nomenclature is used for nitrilotriacetate-containing species. When the degree of protonation is well defined, Y represents N(CH3C02)3, e.g. Fey-.When the degree of protonation is not explicitly indicated, NTA is used, as in -Fe"'-NTA or [NTA]. Bonds prior to Fe indicate surface species.M. del V . Hidalgo et al. 11 Results and Discussion Dependence of Dissolution Rate on [NTA] and [Fe"] Fig. 1 shows a series of typical dissolution profiles, in the form of the volume of acid consumed against time; the initial dissolution rates dV/dt are obtained from them and transformed according to eqn ( 1). Generally, the profiles are monotonously deceleratory, the initial rate being the maximum rate. In certain cases an initial dead time was observed, which was never longer than 2 min; it was attributed to initial equilibration processes and eliminated in the calculation procedures. The variation of the initial dissolution rate as a function of the total concentration of nitrilotriacetic acid ([NTA]) at constant total concentration of iron(1r) ([Fe"]) is shown in fig.2. The dependence of the initial dissolution rate on [Fe"] at constant [NTA] is shown in The results in fig. 2 may provisionally be interpreted in terms of one or both of the fig. 3. following complexation equilibria K s [the actual charges involved may be different, see ref. (24), and surface species will be represented below as -Fe-NTA] OH + Y3- + -Fe"I-Y2- + OH- (3) -FeIII- K e The results in fig. 3 may be interpreted qualitatively by considering that the reductant involved is the Fey- ion, and that the process is first-order in this ion. The curvature of the plot of R us. [FeI'] is in this interpretation a consequence of the increase in the uncomplexed Fe" fraction when [Fe"] is increased.* If the interpretation of fig. 3 in terms of equilibrium (4) is correct, the results in fig. 2 should also reflect the complexation of Fe" in solution [eqn (4)]. Attempts to interpret the data in fig. 2 on the basis of equilibrium (4) alone fail: it is not possible to obtain a linear plot of R as a function of [Fey-] from fig. 2 using tabulated K , values.23 Even though deviations from the tabulated values due to medium effects could in principle account for the discrepancy,? a more straightforward interpretation can be obtained by also taking into account equilibrium (3). There is evidence from related systems in this sense: maxima are observed in the plots of R us. [L] at constant [Fe"] in related systems (e.g.L = EDTA) and this is to be ascribed to a retarding effect of equilibria such as (3)." By analogy, we must consider the possibility of an outer-sphere reaction of Fey- with complexed surface sites -Fe'"-NTA and by an inner-sphere reaction with uncomplexed surface sites -Fe"'-OH, the latter process having the larger rate constant : -Fe"'-NTA + Fey- --+ dissolution ( 5 ) Fe(H20)i'+Y3-+FeY-+6H,0. (4) k,. 0 The maximum observed in the Fe30,-EDTA system2' is due to an adequately high value of the ratio kl.o/kl.l and a relatively low value of the ratio K,/K,, the latter * From fig. 3, it is not possible to rule out a small contribution from a reaction pathway independent of Fe". Blank experiments show that in longer time spans, magnetite dissolves to some extent in Fell-free NTA solutions.In the absence of NTA, no dissolution takes place even in the presence of substantial Fe" concentrations [cf: however ref. (17)]. We have analysed the equivalent reaction in EDTA media and a reductive pathway has been suggested, the EDTA itself being the reductant.'.* t The introduction of ionic strength corrections causes only slight changes in the R us. [Fey-] profiles. Even though this is not conclusive evidence, it also points to the line of reasoning given in the text.12 3 m 6 a 2 2 1 Dissolution of Magnetite by Nitrilotriacetatoferrate(n) I I I I I 20 40 60 80 100 t/min Fig. 1. Volume of HCl(O.1080 mol dmP3) consumed to maintain constant pH as a function of time. Experimental conditions: 303 K, pH 2.8, magnetite mass 17.9 mg, solution volume 50 ~ m - ~ , 102[NTA] and 1O2[FeI1] (moldm-3) as follows: 0.60, 2.41 (a); 1.86, 2.12 (0); 3.87, 2.10 (A); 4.16, 2.91 (B); 4.19, 4.19 (A).4.0 c( I ," 3.0 E E 'p w - s 2.0 1 .o 0.5 1.0 2.5 3.0 [NTA]/lO-' mol dme3 Fig. 2. Initial dissolution rates R as a function of total [NTA] at [FeT1] = 2.1 x mol dm-3. Other conditions as in fig. 1.M. del V. Hidalgo et al. 13 I I I I 1.0 2.0 3.0 L.0 [ Fe"]/ 1 0-2 mol dm-3 Fig. 3. Initial dissolution rates R as a function of total [Fe"] at WTA] = 4.17 x mol dm-3. Other conditions as in fig. 1. defining the possibility of separating volume complexation of Fe" [eqn (4)] from surface coverage [eqn (3)]. In the present system, KJK, is considerably higher,? and this could account for the absence of a maximum.The two parallel paths (5) and (6) could in principle also explain the curvature of the R us. [Fey-] plot that can be derived from fig. 2 using the K, value from the l i t e r a t ~ r e . ~ ~ Fig. 4 shows that R/[FeY-] decreases as WTA] increases, in a fashion suggesting an (inverted) adsorption isotherm (except at very high [NTA], see below); however, the implied value of K, is too low to be reasonable, and an alternative explanation must be sought. According to the usual ideas about electron-transfer reactions,25 the rate constants kl,o and k l , l cover a complex sequence of events. As stated above, there are no experimental conditions under which pathway (5) becomes negligible. On the other hand, complete surface coverage by NTA is achieved in many of our experiments, and it is therefore possible to analyse kl,o in more detail. Process (5) in fact represents scheme 1.K* - FeIII-NTA + Fey- a- - FeIII-NTA Fey- (precursor complex) ket 11 k-et products - kdis - FeII-NTA. FeY (postcursor complex) Scheme 1. t This assertion is borne out by an expe.rimenta1 study of the adsorption of NTA, which shows that K,(NTA) is not much smaller than KJEDTA); on the other hand, Ke(NTA) is smaller than Ke(EDTA).2314 Dissolution of Magnetite by Nitrilotriacetatoferrate(11) 0.05 0.10 ' [ NTA]/mol dm3 - 0.30 Fig. 4. Changes in the ratio R/[FeY-] as a function of WTA]. Experiments are those presented in fig. 2. [NTA] has been corrected for complex formation in solution. The rate will be first-order in [Fey-] only if K*[FeY-] -g 1, i.e. if precursor complex formation is an unfavourable equilibrium, and if also kdis 9 kwet.Even without removing the latter condition, we may expect deviations from the first order if K*[FeY-] approaches unity. In that case eqn (7) follows P[ Fey-] 1 + P[FeY-] R = ket[-Felll-NTA] (7) and fits the data in fig. 4 reasonably well if K* = 2.5 x lo3 dm3 mol-'. This is a rather high figure, but not unreasonable in view of the possible interactions of Fey- with surface-bound NTA species, especially at high NTA concentrations. At [NTA] = 0.28 mol dm-3, the rate is higher than would be predicted from the above ideas. Several possible explanations can be put forward (conformational changes in surface -FeIII- NTA species, participation of Fey:-, medium effects, etc.) but none adequately substantiated.Gorichev and Kipriyanov2' have analysed the consequences of k-et 9 kdis; for this case they conclude that log R should vary linearly with log [a(Fe")/a(Fe"I)], a representing activities, the slope being 0.5. In the series of experiments included in fig. 2, the changes could then be attributed to the shifts in the FeIII-FeII couple potential as FJTA] is increased. Plots of log R against log ([Fe'I']/[Fe"]) taking into account the formation of Fey:-, FeY and Fey- are in fact roughly linear, the slope being ca. 0.33. However, addition of FeIII does not influence the rate appreciably, and this rules out a possible control by the FeI' phase-transfer process : even though the actual exponent for Ferrl (-0.5) for this case could in principle be changed without altering the essence of the argument, fast reverse electron transfer should give rise to a negative order on FeIII that is not observed experimentally in this case.7 The above discussion emphasizes the complexities and ambiguities of rate interpretations in these reactions.t At low [NTA], aqueous FerIr scavenges NTA, displacing equilibrium (4) to the left, and total inhibition of dissolution is observed.M . del V. Hidalgo et al. 15 Dependence of the Dissolution Rate on pH Fig. 5 shows the variation of dissolution rate as a function of pH at constant [NTA], [Fe"] and A / V ratio. A decrease in the rate as the pH increases can be seen together with a tendency towards a maximum, which cannot however be observed because H,Y precipitates at lower pH values.This dependence is due to the superposition of various effects :26 (1) the influence of pH on the speciation of complexed ions in solution, (2) the influence of pH on ligand adsorption and (3) the influence of pH on the concentration of active sites. The first factor is related to the changes in [Fey-] with pH due to the variation of [Y"] : WTA]/[Y3-] = 1 + Kii[H+] + Ki:K,-,1[H+l2 + KitKilK;i[H+]3 (8) where the K, are the successive acidity constants of H,Y. Other pH effects can be deduced from this by comparing R/[FeY-] values at various pH. This procedure corrects for changes in [Fey-] due to the relationship described by eqn (8), and the resulting plot given in fig. 6 shows that factors (2) and/or (3) above must be taken into account.Separation of the two remaining factors is qualitatively simple, because of the noted reactivity of both complexed and uncomplexed surface sites. The ratio kl, Jk1, being larger than one, surface complexation should in any case lead to a decrease in rate with decreasing pH; eventually, a minimum rate could characterize the pH of maximum adsorptivity typical of polycarboxylates on iron oxides (pH 2-3).43 '* 27 The influence of surface complexation on the pH dependence of the reaction rate is, however, probably minor; this is especially true for experiments carried out at high NTA concentrations, where surface coverage is large. The data in fig. 6 indicate a rate law of the form R = kH[H+]0.67 (9) for complexed surface sites; k , incorporates the NTA dependence.The exponent 0.67 is typical of dissolution kinetics, and it is generally interpreted as representing a Freundlich- type adsorption equilibrium (see below for a further possibility regarding changes in surface potential). Thus, eqn (9) represents a requirement for protons adjacent to the site where the reductant is adsorbed. The dissolution sites must therefore be visualized as kinks, where the bonding strength between -Fe"' and residual lattice oxide anions is decreased by protonation. Other systems'' 2 5 9 28 exhibit exponents 0.5 that are also in line with this reasoning. An important point to consider in the analysis of pH influence is the pH dependence of the surface potential, ly,, and the potential in the inner Helmholtz plane, lyB. For dissolution reactions controlled by the rate of transfer of cations across the interface, a model has been proposed'' that incorporates into the pH dependence of the rate an equilibrium factor (surface excess of H+) and a kinetic factor derived from the term exp (cczFAly/RT), where a is the electrochemical transfer coefficient, z the algebraic charge number of the transferring ion, F = 96500 C and Aly = y o - lyp This kinetic factor is indeed important in the dissolution of magnetite by sulphuric acid.17 In the present case, rate control by cation transfer is not indicated by the rate data.Iron transfer should take place as Fe'INTA and any increase in yo due to proton adsorption should hinder such a process. On the other hand, electron transfer from Fe'INTA in the inner Helmholtz plane to -Fe'I'-NTA should be assisted by increasing Aly.Aly cannot be easily modelled as a function of pH in polycarboxylate-containing systems as compared to simple mineral acid systems," because a large number of adjustable parameters are involved in the calculations. At high adsorption densities it is likely that surface potential is not very sensitive to pH. At high coverages electron16 -5.5 -6.0 bo 3 -6.5 Dissolution of Magnetite by Nitrilotriacetatoferrate(Ir) ' . 2 5 30 35 40 4 5 Fig. 5. Initial dissolution rates R as a function of solution pH: [Fe"] = 2.07 x lov2 mol dm+, PH [NTA] = 5.96 x mol dm-3. Other conditions as in fig. 1. I I I 5 4 3 PH Fig. 6. Log R* as a function of pH : data from fig. 5. R* is the initial reaction rate R corrected for changing [Fey']; the corrections were performed using the data in fig. 4.transfer itself, being an outer-sphere process, can be envisaged within the Marcus formalism; this is analysed in detail in a forthcoming paper.2o The driving force for the process is then related to a more negative potential for the couple Fe'II-Fe'' in the inner Helmholtz plane as compared to the analogous surface couple. In this context, changes in by/ are probably only minor perturbations that need not be considered.M. del V. Hidalgo et al. 17 - 6 -65 - 7.0 * 5 -75 - 8.0 3.00 310 3.20 3 30 K/ T Fig. 7. Arrhenius plot. Rate constants (s-l) were obtained in these experiments from the slope of [ 1 - (1 -f11/3] vs. time plots; similar results are obtained from initial rates.Experimental conditions: [NTA] = 3.8 x mol dm-3, pH 3.0. Other conditions as in fig. 1. mol dm-3, [Fe"] = 2.2 x Dependence of Dissolution Rate on the Temperature The Arrhenius plot shown in fig. 7 indicates an apparent activation energy of 73 kJ mol-l. According to the reaction scheme outlined above, this value is a composite of equilibrium and kinetic parameters involved in scheme 1. We thank Dr A. E. Regazzoni for helpful discussions and CONICET for partial support ~ - _ _ - . . - - - - _ _ _ _ . and a fellowship (to M. del V. H.). 2 M. A. Blesa and A. J. G. Maroto, in Decon;amination of Nuclear Facilities (Canadian Nuclear 3 J. Rubio and E. Matijevic, J. Colloid Interface Sci., 1979, 68, 408. 4 H. C. Chang, T. W. Healy and E. Matijevic, J. Colloid Interface Sci., 1983, 92, 469.5 H. C. Chang and E. Matijevic, J. Colloid Interface Sci., 1983, 92, 479. 6 H. C. Chang and E. Matijevic, Finn. Chem. Lett., 1982, 90. 7 M. A. Blesa, E. B. Borghi, A. J. G. Maroto and A. E. Regazzoni, J. Colloid Interface Sci., 1984, 98, 8 E. H. Rueda, M. A. Blesa and R. L. Grassi, J. Colloid Interface Sci., 1985, 106, 243. 9 N. Kallay and E. Matijevic, Langmuir, 1985, 1, 195. Association and American Nuclear Society, 1982), p. 1. 295. 10 Y. Zhang, N. Kallay and E. Matijevic, Langmuir, 1985, 1, 201. 11 N. Valverde and C. Wagner, Ber. Bunsenges. Phys. Chem., 1976, 80, 330.18 Dissolution of Magnetite by Nitrilotriacetatoferrate(11) 12 N. Valverde, Ber. Bunsenges. Phys. Chem., 1976, 80, 333. 13 N. Valverde, Ber. Bunsenges. Phys. Chern., 1977, 81, 380. 14 J. W. Diggle, Dissolution of Oxide Phases, in Oxides and Oxide Films, ed. J. W. Diggle (Marcel Dekker, 15 M. G. Segal and R. M. Sellers, J. Chem. SOC., Faraday Trans. I , 1982, 78, 1149. 16 E. C. Baumgartner, M. A. Blesa, H. Marinovich and A. J. G. Maroto, Znorg. Chem., 1982, 22, 2224. 17 V. I. E. Bruyere and M. A. Blesa, J. Electroanal. Chem., 1985, 182, 141. 18 M. G. Segal and R. M. Sellers, Advances in Inorganic and Bioinorganic Mechanisms, ed. A. G. Sykes (Academic Press, London, 1984), vol. 3, p. 97 and references therein. 19 M. A. Blesa, A. J. G. Maroto and P. J. Morando, J. Chem. Soc., Faraday Trans. I , 1986, 82, 2345. 20 M. A. Blesa, H. Marinovich, E. C. Baumgartner and A. J. G. Maroto, Znorg. Chem., in press. 21 E. B. Borghi, A. E. Regazzoni, A. J. G. Maroto and M. A. Blesa, to be published. 22 A. E. Regazzoni, G. A. Urrutia, M. A. Blesa and A. J. G. Maroto, J. Znorg. Nucl. Chem., 1981, 43, 23 A. E. Martell and R. M. Smith, Critical Stability Constants (Plenum Press, New York, 1974), vol. I. 24 A. E. Regazzoni, M. A. Blesa and A. J. G. Maroto, 55th Znt. Congr. Colloid Surf. Sci. (Potsdam, 25 N. Sutin, Prog. Znorg. Chem., 1983, 30, 441. 26 I. G. Gorichev and N. A. Kipriyanov, Russ. J. Phys. Chem. (Engl. Transl.), 1981, 55, 1558. 27 M. A. Blesa and A. J. G. Maroto, in Reactivity of Solids, Materials Science Monographs, ed. P. Barret 28 B. Terry, Hydrometallurgy, 1983, 11, 315. New York, 1973), vol. 2. 1489. New York, June 1985). and L-C. Dufour (Elsevier, Amsterdam, 1985), vol. 28A, p. 529. Paper 611057; Received 28th May, 1986

ISSN:0300-9599    

DOI:10.1039/F19888400009

出版商:RSC    

年代:1988

数据来源: RSC

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J . Chem. SOC., Furuduy Trans. 1, 1988, 84(1), 19-28 The Second Planar Virial Coefficient for Nitrogen, Oxygen and Carbon Monoxide adsorbed on Graphite Leszek tajtar" and Stefan Sokolowski Department of Theoretical Chemistry , Institute of Chemistry, M. Curie- Sklodo wska University, 20-031 Lublin, PI. M. Curie, Sklodowskiej 3, Poland We present the results of calculations of the second planar virial coefficients for N,, 0, and CO adsorbed on graphite. In particular, we discuss the effect connected with the influence of the solid substrate upon interaction energy of a pair of adsorbed molecules, the role of electrostatic interactions and the effects of deviations of adsorbed film from an ideal planar configuration. Numerous experimental and theoretical studies of physical adsorption of small linear molecules on well characterized solid surfaces have indi~atedl-~ that at coverages below one monolayer and at sufficiently low temperatures the adsorbed fluid forms an almost two-dimensional film with molecules oriented parallel to the surface.To a first approximation the thermodynamic description of such systems can be formulated6, ' in terms of a two-dimensional model, by neglecting out-of-plane motion of adsorbed molecules. Different theoretical techniques, such as thermodynamic perturbation theory, can be proposed to calculate the properties of adsorbed two-dimensional submonolayer molecular fluid.6 In the limit of very low surface coverages, the properties of such systems can be represented exactly by the virial expansion.*-'' In addition, a comparison of the theoretically predicted and experimentally evaluated values of the virial coefficients constitutes an unambiguous test for the model potentials used in calculation^.^^ l2 In this work we present the results of calculations of the second planar coefficient for nitrogen, oxygen and carbon oxide adsorbed on graphite.The van der Waals interactions between a pair of adsorbed molecules are represented by the sum of four site-site potentials ; the electrostatic interactions, however, are handled by using the well known analytical formula for an ideal quadrupoleequadrupole energy,13 as well as by modelling these interactions by a discrete molecular charge distribution. 2* 3, '* Moreover, we also consider the effects connected with the substrate screening of the interactions between adsorbed m01ecules,~~-~~ and in the case of N, and 0, the effects connected with fluctuations of adsorbed particles out of a common plane." The Second Virial Coefficient For a diatomic fluid interacting via the potential u(r, O,, O,), the second planar coefficient is given by8 B, = -I JOm r dr J l n f i r , O,, 8,) do, do, 4n where is the Mayer function, r is the separation of the centres of the molecules and 0, and 0, are the angles describing molecular orientations.If the temperature is not too low, some deviations from an ideal two-dimensional configuration of adsorbed fluid can be observed and in such cases one should rather use the so-called pseudo-planar second virial coefficient, which is defined by'' fir, 0',0,) = exp [ - u(r, @,,0,)/kTI - 1 B, = ( K ~ B;- w,) K;, x - W, K;, 1920 Second Planar Virial Coeficient where Bi is the bulk three-dimensional second virial coefficient" and KH is the Henry constant.ASKH = k(R)-l]dR (3) I,,, 2 4 wz = J=2"g(Rl)dRlJ a3 k(R2)- 11JIRl,R2)dR2. where A, is the surface area, g(R) = exp [ - u(R)/kT], u(R) is the gas-solid potential, R is an abbreviation for Cartesian r and orientational o coordinates, and the cluster integral W2 is given by'l (4) In the above the subscript 00 means that the integration is extended over all spac? and z, is the position of the Gibbs dividing surface." The difference between B, and B, can be treated as a measure of deviations of adsorbed layer from a strictly two-dimensional configuration.10* l1 The problem of numerical evaluation of the integrals (3) and (4) can be simplified by considering expansions of angular-dependent functions into spherical harmonics Km(o).18 Introducing the reference frame with the z-axis perpendicular to the solid surface we can define the following expansions : (5) g(R) = ( 4 ~ ) ' C glrn(r) Km(m> Im and f(R1, R2) = 4~ C fil12mlrn2(r1, '2) K,m,(W1) K2m,(w& (6) 11 l2m1 m2 where and I 1 sin 81 sin 02 do1 do2 d41 d 4 2 ~ ~ 1 , ~ 2 ) mjU1) x GmJm2)- (8) Jl= KH = /2>2* kOO(Z)-- 11 dz wz = c n 1 gz10(z1) dz1 Jya3 kl*0(Z2) - 4201 dz2 A, I , m, m2(rl, '2) = Now we consider the simplest case of a flat solid surface. Substituting eqn ( 5 ) and (6) (9) into eqn (3) and (4) we obtain and 4 1 1 2 0 0 ( ~ , Zl, 22) dr.(10) I l l 2 z>z* 121-221 If, additionally, the interparticle potential u(R,, R2) does not depend upon the position of two interacting molecules with respect to the solid surface, the coeffi~ient~llpOO can be related to the spherical harmonic coefficient of the Mayer function calculated in the bulk reference frame, i.e.in the frame with the z-axis passing through molecular centres.'* We have : fi11,00(r12, 21 22) Al1,00(r12, 212) = C f;:,(,m Dkm(aBr) D?-rn(aBr) (1 1) O<m<min(l1,1,) where Dim are the elements of the three-dimensional rotation matrix,ls a = y = 0 and = arccos (z12/r12). In such case the one-dimensional integralsL. Lajtar and S. Sokoiowski 21 appearing in eqn (10) can be computed separately and stored, and consequently the integral ( 10) effectively becomes two-dimensional.In the presence of a surface, the molecular interactions become modified. Usually, the modification of the interparticle potential consists of adding to the original (free gas- phase) potential u an additional term Au. In this work we consider the case when Au does not depend upon molecular orientations. Thus f;;zA;0(r12, Zl,Z,) = exp [- Au(r12, Zl,Z,)l lfE(,zz00(~12, z12) + 4 , o SZ201 - 4 , o dZ2O (12) where the superscripts denote the interparticle potentials and the spherical harmonic coefficients fi,z200 can be again determined by using eqn (1 1). The Interaction Potentials In this section we give detailed information about the potentials used in numerical calculations. Because in the case of carbon monoxide our calculations are restricted to the two-dimensional model we will not present here the CO-graphite potentials.Information concerning the interactions of CO with graphite can be found in ref. (2) and (3). Nitrogen and Oxygen The interactions between two gaseous N, and 0, molecules are described by the sum of four site-site Lennard-Jones (1 2,6) terms to which the quadrupole-quadrupole potential is added (1 3) dR1, R,) = 4~ C [(0/rijY2 - (0/rJ61 + ue1 f, i 3Q2 ue, = - [ 1 - 5(cos2 0; + cos2 0;) + 2(sin 0; sin 6; cos & - 4 cos 0; cos e;), 4r5 - 15 C O S ~ e; C O S ~ e;] (14) where the primed angles are measured in the reference frame with the z axis along the line joining the centres of mass of the two molecules and Q is the quadrupole. The values of the parameters of the potentials (13) and (14) are 9s follows (the subscripts denote the interacting species) : E5N/k = 35 K, cr,, = 3.32 A, Qgz = (QN2/€” aiN)r = 1.06, eoo/k = 54.3 K, o,, = 3.05 A, Q& = -0.1965 and the elongations of both molecules are dN2 = 1.1 A and doz = 1.208 A.1292o The electrostatic interactions between two N, molecules can be also modelled by the sum of nine coulombic interactions between three partial charges.Charges of -8.49 x low2’ C are placed at the Lennard- Jones centres and a charge of 16.98 x In the presence of a solid surface, the free gas-phase potential (13) is modified. Following earlier author^^^-^' we allow for this effect using the molecule-molecule version of McLachlan theory [cf. ref. (3)]. In the case of the two-dimensional model we calculate Au according to the equation C at the bond centre.’, where p = 1 + 4L2/r2 and L is the height of the adsorbed layer above an effective image plane. For graphite” (17) L = (2) - dG/222 Second Planar Virial Coeflcient Fig.1. (a) The second planar B, and the second pseudo-planar l$ virial coefficients for different models of nitrogen adsorbed on graphite, Part (a) was calculated assuming the free gas-phase pair energy between two nitrogen molecules, whereas the curves presented in part (b) were obtained by using the surface-mediated interparticle potential. The meaning of all symbols is explained in the text. (a) (-) 2QG, (---) 20G, (...) 30G, (--) 3QG, 0, 2DG; (b) (-) 2QS, (----) 20S, (--) 3QS, .,2DS. where dG = 3.37 A is the interlayer spacing of graphite and the distance (z) of the adsorbed layer from the top carbon layer can be evaluated from experiments [cf.ref. (16)] or from computer sir nu la ti on^.^ In many cases (2) can be approximated by the value of z at which the molecule-solid potential reaches its minimum value.lS In general (2) depends upon temperature, but in our calculations we assume that (2) is independent of temperature. A suitable generalization of eqn (16) to the three- dimensional case was developed in ref. (3). We have (18) where r;, is the distance between one of the adatoms and the image of the others in the substrate; and r12 and 4, subtend the angles a, and a, with respect to the surface: ri2 = (r;2 + 42, z,>+ a, = arcsin (12, -z21/r12) a, = arcsin [(z, + z2)/ri2]. Au(r,,, z,, 2,) = C,(2 + 3 cos 2a, + 3 cos 2a2)/6(r12 r'12)3 - C2/(r;2)6 In our calculations we seto Cl, N,/k = 23 1 288.26 K kJ C,, N,/k = 1 17 267.804 K ( z ) ~ ~ = 3.1 A.21 l7 Cl,02/k = 225953.8446 K A6, C2,02/k = 97775.9 K A6,17 ( z ) , ~ = 3-32 A5*12 andL.tajtar and S. Sokoiowski potentials. In our calculations we assume that the surface is completely flat and that 23 The molecule-substrate potential was assumed to be the sum of two atom-surface U(Z, 8) = U,(Z + 0.5d cos 8) + V,(Z - 0.5d cos 8) where a, = 2.46 A, oAG = (o+o,)/F, o, = 3.4 A, E,, = ( E , E ~ ) ~ and Ep/k = 15 K.22*23 Thus ENG/k = 31.3 K, oNG = 3.36 A, E,,/k = 39 K and o, = 3.225 A. The periodic terms neglected in eqn (19) are and we believe that their influence on the value of the second pseudo-planar virial coefficient is negligible [cf.ref. (24)]. Carbon Monoxide The van der Waals interactions between two CO molecules are described by the sum of four (exp-6) atom-atom potentials3 (20) with AC9/k = 3.81271 x lo7 K, A,./k = 3.60824 x lo7 K, A,, = 3.90893 K, a,, = 0.239 23 A, a,, = 0.255 75 A, a,, = 0.271 94 A, B,,/k = 21 1.8038 K A6, B,,/k = 170.79001 K A6, B,,/k = 130.4981 K A6 and the molecular elongation was d,, = 1.1282 A.3 The electrostatic interactions are handled by employing a three-site mod$ - 1.018935 x C , 1.231 345 x C and -0.31241 x 1O-l’ C located at - 1.0!2 A, -0.6446 A and 0.3256 A, respectively, with the atomic sites located at -0.6446 A (C) and 0.4836 A (0) with respect to the centre of mass. The effect of substrate screening on the pair intezaction energy is again described by the potential 116) with3 C,,,,/k = 272 181.66 K A6, C2,,,/k = 143 367.45 K A6 and (z),, = 3.35 A.uij(rij) = A exp ( - rij/a) - B/rfj Results and Discussion We begin our characterization by defining abbreviations for the system studied. The numbers 2 and 3 will refer to the planar and pseudo-planar virial coefficient. The symbols Q and D will denote the ideal quadrupole potential (14) and the discrete charge distribution model [eqn (1 5)] used in calculations of electrostatic interactions, whereas the label 0 means that the electrostatic interactions are completely neglected. The symbols G and S denote the free gas-phase potential and the substrate-mediated interactions between a pair of molecules. For example, the symbol N,-2QG means that the calculations for N, are performed according to the two-dimensional model assuming that the quadrupole-quadrupole interactions are described by eqn (14) and that the influence of the underlying solid on the interparticle potential is neglected. The numerical integrations were performed by using standard multidimensional Gaussian and Gaussian-Chebyshev procedures [cf.ref. (25)]. The calculation of the sum (10) was truncated after the term with 1, = 1, = 6. Our previous investigations26 have indicated that for molecular elongations considered here, the higher-order terms in eqn (10) give quite negligible contributions. The dependence of the virial coefficients upon temperature can be described by using the formulae collected in table 1. The proposed approximations recover the values of the two-dimensional virial coefficients with an accuracy of 1 .O O/O, except in the vicinity of the Boyle temperatures (see table l), where larger deviations are observed.In the case of pseudo-planar virial coefficient this accuracy is worse and is 5 % . The variation in the second planar and pseudo-planar virial coefficients with temperature is also presented in fig. 1-3. The following observations can be made. (a) The effect of electrostatic interactions on the values of the virial coefficients computed for oxygen is very small. On24 Second Planar Virial Coeficient Table 1. Analytical approximations describing the temperature dependences of the second planar and pseudo-planar virial coefficients and the Boyle temperatures estimated for the investigated systemsa system Boyle temperature designation a b C d temperature/K range/K N2-20S 173.75 19.46 0.571 2.084 105 40-270 N,-2QS 168.04 22.01 0.7 13 2.140 102 40-280 N,-2DS 171.52 21.89 0.930 2.196 102 40-240 N2-20G 220.76 21.93 1.050 2.228 127 40-240 N,-2QG 215.74 24.12 1.212 2.297 125 40-260 N,-2DG 213.70 24.41 1.191 2.276 125 40-240 Ni-3QS 151.92 13.80 2.040 1.071 N2-30G 318.86 10.74 3.876 2.977 N,-3QG 234.55 23.89 1.836 2.184 O,?") 259.23 26.10 0.78 1 2.371 2QS 02JoG] 320.31 32.12 0.738 2.358 2QG 21 50-250 42 50-250 45 50-250 40 50-240 75 60-270 02-3QS 383.85 13.61 3.469 3.193 156 50-250 0,-3QG 434.63 23.42 3.494 3.101 205 50-250 c o - 2 0 s 254.90 39.92 0.351 2.234 155 70-240 CO-2DS 321.19 66.82 3.333 3.399 183 90-260 CO-2DG 302.13 76.23 0.197 2.298 210 110-240 a The general approximating formula is B, = -a/( T- b) - 0.001cT+ d.the other hand, the electrostatic interactions significantly effect the values of the virial coefficients determined for carbon monoxide. Over the whole range of investigated temperatures the curve evaluated for the CO-2DS model lies below the curve labelled by 20s. (b) In the case of nitrogen both the ideal quadrupole model and the discrete charge distribution model lead to very close values of the virial coefficients. (c) The differences between pseudo-planar and planar virial coefficients computed for nitrogen and oxygen are not negligible, even at low temperatures. Over almost the entire range of temperatures the values of the pseudo-planar virial coefficients are lower than the values of the planar virial coefficients.( d ) Modification of a pair potential by the presence of a solid surface causes a shift of the computed curves in the direction of lower temperatures. (e) We observe different effects of quadrupole-quadrupole interactions on second pseudo-planar and second planar virial coefficients of nitrogen adsorbed on graphite. These interactions reduce the values of the pseudo-planar virial coefficients, and at temperatures >80 K the values of the planar virial coefficients become higher than corresponding values of the virial coefficients evaluated for models which neglect the electrostatic interactions. To explore the reason for such different behaviour of pseudo-planar and planar virial coefficients of nitrogen we evaluated the two- and three-dimensional Boltzmann- averaged potentials (21) and (22) i&)(r) = -kT In (exp [ - u(r, el, B,)/kT])ff,i i&D(r) = - kT In Gfooo + 1) = - kT In (exp [ - u(R1, R,)/kq);D+ where (...)XI$: and (...>rAa denote unweighted averages over two- and three- dimensional rotations.Both potentials ziZD(r) and aaD(t-) depend upon temperature, andL. Lajtar and S. Sokoiowski 25 300 100 200 1 " " 1 " " 1 ' " ' 1 I ' TIK B Fig. 2. Fig. 2. As in fig. 1, but for oxygen. (- TIK 100 200 1 " " 1 " " 1 " " 1 " " 1 " B Fig. 3. ) 2QS, (---) 2QG, (--) 3QS, (....) 3QG. Fig. 3. As in fig. 1, but for carbon monoxide. (-) 2QS, (---) 20S, (--) 2QG. in general fi,,(r) # ~ ~ ~ ( r ) . We note that the definition of the second planar coefficients, eqn (l), can be rewritten as B = -0.5 {exp[-zi,,(r)/kT]- 1)dv (23) I and the leading term in the sum (10) involves the integral of (exp [ - u3,(r)/kT] - l}.In fig. 4 and 5 we give a comparison of the Boltzmann-averaged potentials computed for nitrogen and carbon monoxide. The addition of a quadrupole term to the intermolecular potential of nitrogen causes that the minima in the #3D(') curves to become deeper, and this effect does not depend upon temperature. A similar effect of electrostatic interactions is observed in the case of the potential UID(r) computed for carbon monoxide. The situation is different in the case of the two-dimensional potential #,,,(r) computed for nitrogen. At low temperatures the quadrupole-quadrupole interactions reduce the minimum in uzD(r), but at higher temperatures the addition of the quadrupole term to the nitrogen-nitrogen potential causes this minimum to become shallower. We can thus state that the observed different effects of quadrupole interactions upon the values of planar and pseudo-planar virial coefficients of nitrogen can be attributed to a different averaging of the pair potentials over angles in two and three dimensions.However, in all cases the modification of a pair potential by molecule-surface interactions causes the Boltzmann-averaged potential minima to be less pronounced. The calculations performed in this work clearly demonstrate that the application ofI ' I 3 L 5 1 I I I I 1 I 1 I L 5 6 rlA 3 L 5 I I I I I ( b ) I I I I I L 5 6 rlA I 1 1 I L 5 6 rl A Fig.4. Two- and three-dimensional Boltzmann-averaged potentials for nitrogen adsorbed on graphite. All symbols are explained in the text, (a) (-) 3QG, (---) 30G; (b) (-) 2QG, (---) 20G; (c) (-) 2QS, (---) 20s.L. Eajtar and S. Sokolowski 27 10 5 - i - 1: - 5 - 1c L 5 6 r i K Fig. 5. As in fig. 4, but for carbon monoxide (-) 2DS, (---) 20S, ( . . . ) 2DG at T = 200 K and (---) 20s at T = 100 K. a strictly two-dimensional model to the description of real adsorption systems may in some cases produce unexpected errors. The discrepancies between two- and three- dimensional virial coefficients would be at least partially removed by replacing the two- dimensional potential GzD(r) in eqn (23) by the three-dimensional potential UBD(r). The last hypothesis should be tested by comparing computed virial coefficients with those evaluated experimentally. References 1 R.D. Diehl and S. C. Fain Jr, Surf: Sci., 1983, 125, 116. 2 J. Piper, J. A. Morrison and C. Peters, Mol. Phys., 1984, 53, 1463. 3 C. Peters and M. L. Klein, Mol. Phys., 1985, 54, 895. 4 R. P. Pan, R. D. Etters, K. Kobashi and V. Chandrasekharan, J . Chem. Phys., 1982, 77, 1035. 5 J. Talbot, D. J. Tildesley and W. A. Steele, Faraday Discuss. Chem. Soc., 1985, 80, 119. 6 L. tajtar, J. Penar and S. Sokolowski, J . Chem. Soc., Faraday Trans. I , 1987, 83, 1405. 7 T. Boublik, 1985, Mol. Phys., 1985, 54, 1644. 8 J. S. Rowlinson, J. Talbot and D. J. Tildesley, Mof. Phys., 1985, 54, 1065. 9 W. A. Steele, The Interactions of Gases with Solid Surfaces (Pergamon Press, Oxford, 1974). 10 J. R. Sams, Prog. Surf. Membr. Sci., 1973, 8, 1. 11 S. Sokolowski, J . Chem. Soc., Faraday Trans. 2, 1981, 77, 405. 12 J. Talbot, D. J. Tildesley and W. A. Steele, Mol. Phys., 1984, 51, 1331. 13 J. D. Hirschfelder, C. F. Curtis and R. B. Bird, Molecular Theory of Gases and Liquids (Wiley, New 14 P. A. Monson, W. A. Steele and W. B. Streett, J. Chem. Phys., 1983, 78, 4126. 15 A. D. McLachlan, Mof. Phys., 1964, 7 , 381. York, 1954).28 Second Planar Virial Coeficient 16 S. Rauber, J. R. Klein and M. W. Cole, Phys. Rev. B, 1983, 27, 1314. 17 S. Rauber, J. R. Klein, M. W. Cole and L. W. Brunch, SurJ Sci., 1982, 123, 173. 18 J. R. Sweet and W. A. Steele, J . Chem. Phys., 1967, 47, 3029. 19 M. E. Rose, Elementary Theory of Angular Momentum (Wiley, New York, 1957). 20 C. A. English and J. A. Venables, Proc. R. SOC. London, Ser. A , 1974, 340, 57. 21 L. tajtar and S. Sokolowski, Czech. J . Phys., in press. 22 W. A. Steele, Surf. Sci., 1973, 36, 317. 23 W. A. Steele, J. Phys. Paris (Colloq.), 1977, 38, C4, 61. 24 W. A. Steele, Surf Sci., 1973, 39, 149. 25 F. Lado, Mol. Phys., 1982, 47, 283; 299. 26 S. Sokolowski, Phys. Lett., 1986, 117, 468. Paper 6/1421; Received 17th July, 1986

ISSN:0300-9599    

DOI:10.1039/F19888400019

出版商:RSC    

年代:1988

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans I, 1988, 84(1), 29-36 The Interaction of Water with Non-electrolytes The System Water-Acetonitrile-1.,4-Dioxane Hero Mirti" and Vincenzo Zelano Dipartimento di Chimica Analitica, Universita' di Torino, via Giuria 5 , I10125 Torino, Italy 'H, 13C and 1 7 0 magnetic resonance has been used to investigate the degree of modification of the aqueous structure in water-acetonitrile-l,4-dioxane mixtures of different compositions. Chemical-shift data of the hydrogen and oxygen nuclei of water have enabled the calculation of a parameter cf,), which may be considered representative of the average situation experienced by one molecule of water in the presence of one molecule of non-electrolyte. Chemical-shift data of the hydrogen and carbon nuclei of acetonitrile and dioxan have been used to determine similar parameters uiN and f,"") representative of the situation experienced by one molecule of water in the presence of one molecule of acetonitrile or dioxane, respectively.The values of fiN and f,"" accord with those off,, and indicate that the degree of modification of the aqueous structure is generally independent of the composition of the non-electrolyte. It is known that solvent features can affect chemical processes in solution. The observed effects, such as a change in the reaction rate or a shift of equilibrium conditions, can frequently be correlated with the physico-chemical properties of the medium. As a change in the latter can dramatically alter the chemical behaviour of a given compound, it is often useful to search for the most suitable medium by mixing different solvents in appropriate amounts.In this way some properties can be varied, e.g. the donor capability or the dielectric constant, simply by varying the composition of the medium. Because of its role in the natural environment, water is one of the most recurrent components of mixed solvents. The addition of a cosolvent to water alters the hydrogen- bonded structure of the latter, and this may be one factor in determining the effect of the medium on the behaviour of the species in solution. The most likely modification made to the network of H,O molecules is the distortion or rupture of the existing links,'-' even though some enhancement of the aqueous structure has been observed at low temperatures in the presence of small quantities of certain non-electrolytes.2* 3, '-12 Whether water-water hydrogen bonds are distorted rather than broken in the presence of organic cosolvents can be widely discussed.The most current view seems to favour the idea of a distorted network, but deciding at what point a distorted link is turned into a broken one may be largely a matter of opinion. It seems likely, however, that an actual rupture of links should gradually take place when the amount of cosolvent becomes predominant in the system and fewer molecules of water are dispersed in the organic matrix, with water-water bonds eventually replaced by water-solvent links. One can generally speak of a weakening of the structure of water, as opposed to its strengthening, but it must be borne in mind that the former may be the result of either simple distortion or more severe breaking of the bonds involved.In this context it is useful to remember that the features of the hydrogen bond network have been argued even in pure water,13-18 where the intermolecular links may be distorted (or broken) by a change of temperature. 2930 Interact ion of Water with Non-elec troly tes In a previous paperlg some binary systems formed by water and a non-hydroxylic solvent are reported. The results obtained showed a correlation between the extent of modification of the water structure and the physico-chemical properties of the cosolvent. In this context it was surprising to note that acetonitrile (AN) and 1,4-dioxane (DO) caused very similar effects, in spite of their different properties.A subsequent study of a ternary system of water, dimethyl sulphoxide (DMSO) and DO showed that the non-electrolytes can affect each other's capability of modifying the structure of water.20 This suggested a study of water-acetonitrile4ioxan mixtures, to discover whether AN and DO can replace each other without changing the degree of distortion (or rupture) of the hydrogen-bond network of water. Among the many techniques allowing one to collect information on the structural features of mixed solvents, nuclear magnetic resonance is one of the most suitable, because the chemical shift (6) of the nuclei of the molecules involved is affected by the structural changes. In this work, magnetic resonance of all 'H, 13C and 170 nuclei of water, acetonitrile and dioxan has been used to study the system under investigation.Experimental Dried acetonitrile and 1,4-dioxane (Riedel-De Haen, maximum water content 0.03 and 0.0 1 YO, respectively) were mixed with demineralized, twice-distilled water to prepare samples in which the mole fraction of each component could vary from 0 to 1. For each mixture, 5 x loh3 dm3 were prepared by using an Amel 233 digital automated burette operated by an Apple IIe computer. The precision of the buret (& 1 x dm3) allowed one to obtain standard deviations of the mole fraction of each component in the range 'H n.m.r. spectra were recorded by a Varian T-60 spectrometer using tetramethylsilane (TMS) or sodium 4,4-dimethyl-4-silapentane sulphonate (DSS) as the internal reference for the determination of chemical shifts.The solubility of the reference compounds in the samples under investigation determined the choice between DSS and TMS, but the latter was used whenever possible. A Jeol GX-270/89 spectrometer was used for recording 13C and 170 n.m.r. spectra, and deuterated dimethyl sulphoxide and methanol were used as external references in the determination of the chemical shifts of the carbon and oxygen nuclei, respectively. The instrumental resolution was as much as 0.01 ppm for 'H and 13C, and 0.05 ppm for 170. This led to kO.01 error in the values of a and k0.02 in those off, (see below for definitions). Exceptions were a values obtained from proton resonance of AN and DO (error up to k0.08) and fM values in the water-rich region (error up to kO.1 for xw > 0.9).Results and Discussion (1-5) x 10-3. The chemical shift of the water protons in a system containing a network of more or less distorted hydrogen bonds is given by 6 = C X i 6 , where di is the chemical shift of the protons engaged in hydrogen bonds with a determined degree of distortion and xi the corresponding fractional population with respect to the total number of protons. In the most simple situation, when hydrogen bonds are only broken and not distorted, eqn (1) becomes 6 = x,s;,+x,s; where 8; and 6; are the chemical shifts of the protons which are not hydrogen bonded and normally hydrogen bonded (i.e. without distortion), respectively, and x, and x, areP. Mirti and V. Zelano 31 the corresponding fractional populations.In the most complex situation a continuous distribution of hydrogen bond energies is present, and the sum in eqn (1) must be replaced by an integral. A determination of the true value of 6; is not simple because even pure water can contain broken or distorted bonds, and the energy distribution of these is dependent on temperature.l3-l8 Therefore 6; is more likely to be an unknown value at lower field than that obtainable by n.m.r. measurement on pure water at a given temperature (dN). On the other hand, 6; should be referred to a standard situation where the water molecules can be considered truly free, e.g. as at infinite dilution in a non-interacting solvent or in the vapour state. When water is in the presence of an interacting cosolvent a value (6,) is obtained at lower field than Sl,, owing to the coordinating properties of the cosolvent towards water.Different values'of 6, have been obtained for water molecules diluted in different non-electrolytes, and a correlation has been found between these values and the donor numbers of the cos01vents.~~ The chemical shift of the water protons at a given temperature varies with the composition of a system containing water and one (or more) non-electrolyte(s), moving upfield from the value obtained in pure water at that temperature (6,) to the value extrapolated for water infinitely diluted in that (or those) non-electrolyte(s) (aF). In spite of the complexity of the system, it seems possible to use the experimental chemical shift obtained for a given mixture as a measure of the modifications caused in the structure of water by the presence of that (or those) non-electrolyte(s) at that temperature relative to the two boundary situations represented by 6, and 6,.Because 6, and 6, are different from 6; and 6;, respectively, and depend upon temperature and cosolvent, one can only obtain relative information; however, this can allow one to infer the extent of modification of the water network in the presence of a given non-electrolyte and to compare the modifications caused by different non-electrolytes. In order to attempt a quantitative estimation of the above, one can introduce a parameter (f), which is required as a measure of the average degree of distortion of the water structure provoked by one molecule of non-electrolyte.If different molecules give additive contributions, and if x, is the mole fraction of water in the system, fT1- x,) is a measure of the overall variation of the aqueous structure from pure water at a working temperature. If hydrogen bonds are not distorted but actually broken, f becomes the number of links broken by a single molecule of non-electrolyte. A more convenient way to cope with the matter may be the use of a correlated parameter fM =f/xw, which is related to an average situation experienced by a single molecule of water. A practical reason for usingf, instead offlies in the fact thatfmust be zero for both x, = 1 and x, -+ 0 (i.e. where no cosolvent or water is present, and therefore no structural modifications are possible). On the other hand,f, can range from zero (x, = 1, where the structure is that of pure water at the working temperature) to 2 (x, -+ 0, where all the water-water bonds must be broken and partially replaced by water-cosolvent interactions).As long as hydrogen bonds are only broken and not distorted, eqn (1) is replaced by eqn (2), and this can be rewritten as where fT1 -xw)/2x, is the fraction of water protons not engaged in water-water hydrogen bonds and [2x, -A1 -xw)]/2x, is the fraction of bonded 0nes.l' In this case fM can be obtained as (4) In a more general case, when hydrogen bonds are distorted rather than broken, it seems possible that values off, obtained from eqn (4) may equally well be used to give a picture of the situation created by the presence of a non-electrolyte in water.In fact, because the same experimental value of 6 can stem from only a few broken links f M = 2(SN - 6)/[(6N - 'F) ( l - xw)l' 2 FAR 132 Interaction of Water with Non-electrolytes xw Fig. 1. Variation of the chemical shift of the 'H (a) and I7O nuclei (b) of water as a function of the mole fraction of H,O in mixtures H,O-AN-DO containing equal mole fractions of acetonitrile and dioxane. 3 0 0.2 0.4 0.6 0.8 1 XAN/@AN + XDO) Fig. 2. Variation of the chemical shift of the 'H (a) and I7O (b) nuclei of water as a function of the composition of the non-electrolyte in H,O-AN-DO mixtures containing 0.5 mole fraction of H,O. as well as from many distorted ones, it seems possible to speak in terms of an apparent number of bonds broken by the non-electrolyte.Within the limits of the above assumptions, eqn (4) has been used to calculate values off, from the chemical shifts of both 'H and ''0 nuclei of the water molecules; the values obtained from the hydrogen and oxygen nuclei agree with each other. Fig. 1 shows the variation of the chemical shift of the lH and l 7 0 nuclei with the mole fraction of water in samples containing a fixed molar ratio AN/DO, whereas fig. 2 reports the variation of these chemical shifts as a function of the composition of the cosolvent in samples with a fixed water content. The complete set of proton chemical shifts is given in table 1. The values off, obtained are given in table 2, together with those calculated previously for the binary mixtures H,O-AN and H2O-DO.l9 It is inferred from the data of table 2 that the degree of modification of the aqueous structure is generally independent of the composition of the non-electrolyte except whenP.Mirti and V. Zelano 33 Table 1. Chemical shifts of the hydrogen nuclei of water (aw), acetonitrile (aAN) and dioxane (ano) in H,GAN-DO mixtures of different compositions 0.90 0.90 0.90 0.90 0.90 0.70 0.70 0.70 0.70 0.70 0.50 0.50 0.50 0.50 0.50 0.30 0.30 0.30 0.30 0.30 0.10 0.10 0.10 0.10 0.10 0.90 0.70 0.50 0.30 0.10 0.90 0.70 0.50 0.30 0.10 0.90 0.70 0.50 0.30 0.10 0.90 0.70 0.50 0.30 0.10 0.90 0.70 0.50 0.30 0.10 4.49 4.50 4.51 4.5 1 4.52 4.1 1 4.12 4.15 4.14 4.15 3.76 3.77 3.78 3.79 3.79 3.26 3.3 1 3.34 3.35 3.39 2.58 2.68 2.71 2.74 2.75 3.72 3.72 3.72 3.72 3.72 3.67 3.67 3.68 3.65 3.67 3.66 3.65 3.65 3.64 3.62 3.63 3.62 3.62 3.61 3.61 3.62 3.60 3.60 3.58 3.57 2.06 2.06 2.06 2.06 2.06 2.03 2.03 2.03 2.04 2.03 2.0 1 2.01 2.00 2.03 2.0 1 2.00 2.00 1.98 1.99 2.00 1.98 1.97 1.98 1.96 1.95 0.05 0.05 0.05 0.05 0.05 0.03 0.03 0.03 0.03 0.03 0.02 0.02 0.02 0.02 0.02 0.01 0.0 1 0.01 0.0 1 0.01 0.90 0.70 0.50 0.30 0.10 0.90 0.70 0.50 0.30 0.10 0.90 0.70 0.50 0.30 0.10 0.90 0.70 0.50 0.30 0.10 2.4 1 2.49 2.53 2.55 2.56 2.34 2.40 2.46 2.49 2.50 2.30 2.40 2.45 2.45 2.45 2.22 2.33 2.40 2.43 2.41 3.62 3.60 3.59 3.58 3.58 3.60 3.60 3.60 3.59 3.59 3.60 3.60 3.59 3.59 3.57 3.59 3.60 3.58 3.60 3.58 1.98 1.96 1.96 1.95 1.95 1.97 1.96 1.96 1.96 1.95 1.97 1.97 1.95 1.96 1.94 1.95 1.96 1.95 1.96 1.94 Table 2.Values off, for water-acetonitrile-dioxane mixtures of different compositions X , , / ( X , ~ + X , ~ ) = 0.00 0.10 0.30 0.50 0.70 0.90 1 .oo ~ ~ ~ _ _ _ _ _ _ _ _ ~ ~ ~ ~ 0.10 1.48 1.85 1.82 1.79 1.80 1.71 1.45 0.30 1.59 1.63 1.65 1.62 1.65 1.62 1.57 0.50 1.59 1.54 1.57 1.60 1.59 1.59 1.59 0.70 1.62 1.68 1.66 1.68 I .67 1.62 I .62 0.90 1.84 1.91 1.87 1.90 1.88 1.87 1.84 water is present in a great excess, when one finds an increase infM on passing from AN or DO alone to a mixed non-electrolyte; this means that the capabilities of the two solvents to modify the energy distribution of the water-water hydrogen bonds can be mutually enhanced under certain conditions.This is a situation already in evidence for systems containing water in the presence of both DMSO and DO, and it can be explained if the non-electrolytes interact differently with water, so that a cooperative action can stem from their mixing.This may happen, for example, if the donor capability of one compound (e.g. dioxane) can gain efficacy from the larger dielectric constant of the other (e.g. acetonitrile). 2-234 Interaction of Water with Non-electrolytes B 120 - h E ( o f 118 - W I 67 - 1,,,,,,,,,1 0 0.2 0.4 0.6 0.8 1 XW Fig. 3. Variation of the chemical shift of lH (A) and 13C nuclei (B) of acetonitrile (b) and (a) dioxane as a function of the mole fraction of water in H,O-AN-DO mixtures containing equal mole fractions of DO and AN. 0 0.2 0.4 0.6 0.8 1 0 0.2 0.4 0.6 0.8 1 XAN/(XAN + XDO) XAN/(XAN + XDO) Fig. 4. Variation of the chemical shift of 'H (A) and I3C nuclei (B) of dioxane (a) and acetonitrile (b) as a function of the composition of the non-electrolyte in H,O-AN-DO mixtures containing 0.5 mole fraction of water.Similarly to water nuclei, the chemical shifts of the nuclei of acetonitrile and dixoane (figs. 3 and 4 and table 1) can be considered representative of the variation of the structural features of AN and DO, respectively, when going from the pure cosolvent towards infinite dilution in water through the formation of complex hydro-organic species. In line with the points discussed above, if only two boundary forms contribute to the observed chemical shift, one may calculate a fractional population for each of them. As already said, this may not be the case with water, or indeed with AN or DO, either. However, pure AN and pure DO molecules are not part of a network of strong intermolecular interactions.Therefore, a wide span of intermediate situations is less probable for AN or DO than for water, so that the picture obtained by considering only two boundary structures (that of the pure cosolvent and that of the cosolvent at infinite dilution in water) may be less approximate. In this case one can calculate the fractionalP. Mirti and V. Zelano 35 Table 3. Fractional populations of acetonitrile (a,"") and dioxane (01:~) molecules coordinated to water, calculated from 'H and 13C n.m.r. data of mixtures containing equal mole fractions of either non-electrolyte atN a,Do 'H n.m.r. 13C n.m.r. 'H n.m.r. 13C n.m.r. x,, + X*N 0.10 0.85 0.83 0.8 1 0.86 0.30 0.62 0.58 0.56 0.54 0.50 0.39 0.39 0.38 0.37 0.70 0.23 0.22 0.19 0.20 0.90 - 0.08 0.06 0.1 1 Table 4.Comparison between values off,, xW(x,, + xDo) (from n.m.r. data of water nuclei) and fiNxWxAN + fgoxwxno (from n.m.r. data of acetonitrile and dioxane nuclei) obtained for H,O-AN-DO mixtures containing equal mole fractions of AN and DO 0.10 0.16 0.30 0.34 0.50 0.40 0.70 0.35 0.90 0.17 0.17 0.34 0.38 0.29 0.17 population of acetonitrile and dioxane molecules bound to water (atN and a!*, respectively) from the experimental chemical shifts as where a, may be either a,"" or a:", and 6, and 6, are the chemical shifts of the solvent considered in its boundary situations. By doing so, one neglects possible interactions between AN and DO in the mixed non- electrolyte, but the chemical shifts of the hydrogen and carbon nuclei of AN and DO vary only slightly with the composition of the cosolvent.Such a variation is increasingly less evident as the content of water increases, and disappears completely in an excess of water. This indicates that, even though interactions are possible between AN and DO, they are of much less importance than those of water with each cosolvent. The values of at" and a:" obtained according to the above considerations from data for either protons or carbon nuclei correspond well and are given in table 3. However, values of a,"" could not be obtained from the chemical shift of the oxygen nuclei of dioxane, and the methyl carbon of acetonitrile proved to be of no use in determining a,"". Within the limits of the approximations discussed above, the values of atN and a:" can be used to calculate the mole fraction of acetonitrile and dioxane bound to water (as atNxAN and a,""xD0, respectively, if x A N and x,, are the mole fractions of AN and DO in the system).Those, in turn, can lead one to calculate the mole fraction of the molecules of water bound to each cosolvent, provided that the kind of complex species formed is known. In this respect, a prevailing opinion is that the species formed most easily are those with a 1 : 1 molar ratio of water to cosolvent, even though species with 2: 1 and 1 : 2 molar ratios can be present in a great excess of water or cosolvent, re~pectively.~* ''9 21-24 If one assumes that 1 : 1 species are formed, a,ANxAN and a~"xDO are36 Interact ion of Water with Non - elec tr oly t es also likely to give the mole fractions of water molecules linked to acetonitrile and dioxan, respectively, and these must be multiplied by 2 to represent the mole fractions of water protons.If these two terms are divided by x,, and xDO, separately, and then both of them by x,, one obtains two factors which are representative of the average situation experienced by one molecule of water in the presence of one molecule of acetonitrile and dioxane, respectively (i.e. fkN = 2a,AN/xw and f go = 2a,D0/xw). As long as the overall modification of the structure of water in the presence of both acetonitrile and dioxane stems from additive contributions from the two cosolvents (as shown previously, except for the water-rich region), one can write taking into account that all fM,fiN a n d g o relate to a single molecule of water and non- electrolyte.The agreement between the two halves of eqn (6) (table 4) is particularly remarkable if one considers that the values off,” andgo, on one hand, and those of fM, on the other, are obtained from completely independent measurements of the chemical shifts of cosolvent and water nuclei, respectively. In conclusion, the results obtained confirm that acetonitrile and dioxane lead to a similar average modification of the structure of water and that a variation of the composition of an AN-DO mixed non-electrolyte has no particular effect on that structure. This, of course, does not settle the question as to whether bonds are broken or distorted : however, since the physico-chemical properties of the two non-electrolytes are different, it shows that the same apparent number of bonds broken can stem from water-cosolvent interactions of different kind.References 1 Yu. I. Naberukhin and S. I. Shuiskii, Zh. Strukt. Khim., 1967, 8, 606. 2 C. J. Clemett, J. Chem. SOC. A, 1969, 455. 3 B. Z. Gorbunov and Yu. I. Naberukhin, Zh. Strukt. Khim., 1972, 13, 20. 4 M. F. Fox and K. P. Whittingham, J. Chem. SOC., Faraday Trans. I , 1975, 71, 1407. 5 Yu. A. Volokhov, N. G. Dovbysh, V. B. Lebedev and V. E. Mironov, Zh. Strukt. Khim., 1975, 16, 6 C. Moreau and G. Douhiret, J. Chem. Thermodyn., 1976, 8, 403. 7 A. J. Easteal, Aust. J. Chem., 1979, 32, 1379. 8 D. N. Glew, H. D. Mak and N. S. Rath, J. Chem. SOC., Chem. Commun., 1968, 264. 9 C. J. Clemett, J. Chem. SOC. A , 1969, 458. 1013. 10 T. Tokuhiro, L. Menafra and H. H. Smart, J . Chem. Phys., 1974, 61, 2275. 11 0. D. Bonner and Y. S. Choi, J. Phys. Chem., 1974, 78, 1727. 12 B. Z. Gorbunov, V. S. Kozlov and Yu. I. Naberukhin, Zh. Strukt. Khim., 1975, 16, 808. 13 K. Buijs and G. R. Choppin, J. Chem. Phys., 1963, 39, 2035. 14 D. P. Stevenson, J. Phys. Chem., 1965, 69, 2145. 15 M. R. Thomas, H. A. Scheraga and E. Schrier, J. Phys. Chem., 1965, 69, 3722. 16 0. D. Bonner and Y. S. Choi, J. Phys. Chem., 1974, 78, 1723. 17 M. C. R. Symons, J. M. Harvey and S. E. Jackson, J. Chem. SOC., Faraday Trans. I , 1980, 76, 256. 18 S. A. Rice and M. G. Sceats, J . Phys. Chem., 1981,85, 1108. 19 V. Zelano and P. Mirti, 2. Phys. Chem. (Munich), 1983, 138, 31. 20 V. Zelano and P. Mirti, 2. Phys. Chem. (Leipig), 1986, 267, 857. 21 A. Fratiello and D. C. Douglass, J . Mol. Spectrosc., 1963, 11, 465. 22 A. Le Narvor, E. Gentric and P. Saumagne, Can. J. Chem., 1971,49, 1933. 23 B. Kingston and M. C. R. Symons, J. Chem. Soc., Faraday Trans. 2, 1973, 69, 978. 24 S. 0. Paul and T. A. Ford, Spectrochim. Acta, Part A, 1981, 37, 415. Paper 6/ 1598 ; Received 4th August, 1986

ISSN:0300-9599    

DOI:10.1039/F19888400029

出版商:RSC    

年代:1988

数据来源: RSC

摘要:

J . Chern. Soc., Furuduy Trans. I, 1988, 84(1), 37-40 Effect of the Surface Structure of Metal Oxides on their Adsorption Properties K. Hadjiivanov* and D. Klissurski Institute of General and Inorganic Chemistry, Bulgarian Academy of Sciences, Sofia 1040, Bulgaria A. Davydov Institute of Catalysis, Siberian Division of the USSR Academy of Sciences, Novosibirsk 630090, U.S.S.R. The effect of the first and second coordination spheres on the electron acceptor properties of coordinatively unsaturated metal ions at oxide surfaces has been studied theoretically. Metal ions with the same coordination number can differ strongly in their electrophilic properties depending on the coordination state of the ligands. It is shown that cations at crystal edges can have stronger, equal or weaker electrophilic properties than the corresponding cations situated on the planes forming the edges.The properties of coordinatively unsaturated oxygen ions and surface hydroxyl groups as well as the localization of the strongest Lewis-acid and Lewis-base sites are discussed. It has been e~tablishedl-~ that the Lewis acid sites on a metal oxide surface are coordinatively unsaturated (c.u.s.) metal ions. The variety of these sites for one oxide is due to the variety of surroundings of the metal ions, which are usually classified on the basis of their coordination number with respect to the lattice oxygen. The abundant experimental data3-* on oxide systems show that the number of C.U.S. cations on metal- oxide surfaces is usually much larger than the number of Lewis-acid sites.In addition, there are several types of surface compounds bonded to the metal ions (surface hydroxyl groups, non-stoichiometric oxygen e t ~ . ) . ' - ~ ? 9-13 In these cases the variety of metal-ion properties cannot be attributed to the difference in their coordination number only. The purpose of the present paper is to estimate the electrophilic properties of metal cations with the same coordination number but with different coordination states of the ligands and to summarize the resulting consequences. With most transition-metal oxides the metal-oxygen bond is, to a high degree, covalent. l4 This means that the second coordination sphere, i.e. the coordination state of ligands, may substantially affect the electron-acceptor properties of the central metal ion.Proceeding from these factors, we shall denote the ion state as MR+(al- a,-. . .-a,), where the number of bracketed figures gives the coordination number, the figure itself indicating the vacancies of each ligand, e.g. Mg2+(1-l-l-l-O) denotes a pentacoordinated magnesium ion where four of the oxygen ligands have one vacancy each, and one is coordinatively saturated. Consider a surface complex consisting of an acid-base pair, i.e. a C.U.S. metal ion and a C.U.S. oxygen ion having one vacancy each [fig. l ( a ) ] . Fig. l ( b ) presents the coordination of the oxygen ion from a Lewis acid and the resulting redistribution of electron density. Obviously, the coordination leads to a decrease in electron density around the central metal ion and a corresponding increase in the electron-acceptor properties of this ion, the coordination number remaining the same.A similar coordination from carbon dioxide15 or sulphur trioxide16 is accompanied by a 3738 Effect of Surface Structure on Adsorption Properties Fig. 1. (a) A complex of a C.U.S. metal ion and a C.U.S. oxygen ion with one vacancy each; (b) the same complex after coordination of the oxygen ion with a Lewis acid and the corresponding change m electron density. 0, Metal ion; 0, oxygen ion; 8, Lewis acid. Fig. 2. Scheme of some planes and edges on the titanium dioxide surface. (a) Planes and edges containing the titanium ion in the Ti4+(l-1-0-0-0) state (see text): 1, rutile (101) plane; 2, anatase (001) plane; 3, anatase (101) x (011) edge. (b) Planes containing the.titanium ion in the Ti4+(l-0-0-0-0) state: 1, anatase (100) plane; 2, anatase (101) and (011) planes.(c) Rutile (1 10) plane containing the titanium ion in the Ti*+(O-O-O-O-O) state. a, Pentacoordinated Ti4+, 0, twofold-coordinated 02-, @, tricoordinated 02-. pronounced shift (by ca. 15-30 cm-l) to higher frequencies of the carbon-oxygen stretching modes in the M-CO surface complex. If a C.U.S. metal ion, which is a strong Lewis acid, participates in the coordination, an analogous effect would be observed. The states shown in fig. 1 correspond to metal ions with the same coordination number but situated on different planes (edges, corners). A typical example exhibiting several types of states is titanium dioxide. Fig. 2 shows the structures of some planes and edges characteristic of the surface of anatase and rutile, the states of the titanium ion being Ti4+( I--1-0-0-0), Ti4+( 1-0-0-0-0) and Ti4+(O-0-0-0-0) [fig.2 (a), (b) and (c), respectively]. According to the scheme in fig. 1, transition from the titanium state in fig. 2(a) to the state in fig. 2(b) and, subsequently, to that in fig. 2(c) can occur, the electron-acceptor properties of the titanium ion increasing in the same direction. Thus, the pentacoordinated metal ions on an oxide surface can be ordered in the sequence of increasing electron-acceptor (acidic) properties as follows : (1) Mn+(l--1-1-1-0) or Mn+(l-1-1-0-0), a metal ion in an acid-base network, (2) M"+(l-1-0-0-0), a metal ion from an acid-base row, (3) M"+(l--0-0-0-0), a metal ion participating in an acid-base pair, (4) M"+(O-0-0-0-0), a metal ion with no C.U.S.oxygen ions in its first coordination sphere. One or several of these states can exist on the surface of an oxide crystal. For instance, the pentacoordinated magnesium ions in the (001) plane of magnesium oxide are in theK. Hadjiivanov, D. Klissurski and A . Davydov 39 Mg2+( 1-1-1-1-0) state [excepting Mg2+(2-1-1-1-0) ions situated close to the crystal edges], whereas all four states can be found on the titanium dioxide surface. All ions having a definite coordination number can be ordered in such a way according to their electrophilic properties. The C.U.S. oxygen ions on the surface, which represent Lewis bases, can also be systematized in this way, e.g. the basic properties of the twofold-coordinated oxygen ions of titanium dioxide will increase from fig.2(a) to As a result of the effect of the second coordination sphere (1) hydroxyl groups bonded to metal ion(s) with the same coordination number can differ in properties and spectral registration depending on the state of the metal ion(s) and (2) the strongest acid and base sites of the surface can be situated on different planes (edges, corners). Let us consider two planes of an oxide crystal, one of them consisting of acid-base fig. 2(c). pairs : M"+( 1--0-0-0-0) - 02-( 1-0-0) the other containing a complex of a metal ion with two vacancies and two oxygen ions with one vacancy: O2-(2-O-0)-Mn+( 1-1-0-0)-02-(2-0-0). The stronger Lewis base will be characteristic of the first crystal plane, while the stronger Lewis acid will correspond to the second plane.(3) Metal ions at edges can have the same coordination number but weaker electrophilic properties than the corresponding ions on the planes forming the edge. This can be seen in fig. 2. Titanium ions on the (101) and (01 1) anatase planes have stronger electrophilic properties than those situated on the (101) x (01 1) edges. In view of the fact that the Lewis acid sites in the case of anatase amount to ca. 8 YO of the C.U.S. titanium ions,3 the cations from acid-base rows have to be considered to be inactive since they possess the weakest electrophilic properties. Hence, with decreasing particle size (increase of cation concentration at the edges) the number of acid sites per unit surface will decrease, in agreement with the experimental data presented in ref.(5) and (17). In the literature the properties of disperse oxides are often described by one crystal plane alone, i.e. that predominantly exposed on the surface.** '-137 On the basis of general consideration one may expect less exposed planes to have a higher activity for adsorption due to their higher surface tension.lg To elucidate the surface properties one has to discuss the structure of all forms exposed on the oxide surface to determine the state of ions and to order them according to the decrease in acidic (basic) properties. In high-dispersity systems, where ions situated on crystal edges and corners constitute several per cent of the C.U.S. ions,1'20 the structures of these forms should also be discussed irrespective of the fact that some of them may be inactive towards adsorption.We thank Professor A. Andreev for helpful discussion. References 1 A. Zecchina, S. Coluccia and C. Morterra, Appl. Spectrosc. Rev., 1985, 21, 259. 2 A. V. Kiselev and V. I. Lygin, Infrared Spectra of Surface Compounds (Nauka, Moscow, 1972). 3 A. A. Davydov, IR Spectroscopy Applied to Surface Chemistry of Oxides (Nauka, Novosibirsk, 4 G. Munuera, F. Moreno and J. A. Prieto, Z . Phys. Chem., 1972, 78, 113. 5 K. Hadjiivanov, A. Davydov and D. Klissurski, Kinet. Katal., in press. 6 K. Tanabe, Muter. Chem. Phys., 1985, 3-4, 347. 7 E. Garrone and F. S. Stone, Proc. VZZZth. Int. Congr. Catal., Berlin, 1984 (Verlag-Chemie, Weinheim, 8 R. J. Cvetanovid and Y. Amenomiya, Adu. Catal., 1967, 17, 103. 9 M. Primet, P. Pichat and M-V. Mathieu, J. Phys. Chem., 1971, 75, 1216. 10 M. Primet, P. Pichat and M-V. Mathieu, J. Phys. Chem., 1971, 75, 1221. 11 A. Zecchina, S. Coluccia, E. Guglielminotti and G. Ghiotti, J . Phys. Chem., 1971, 75, 2774. 1984). 1984), vol. 3, p. 441.40 Efect of Surface Structure on Adsorption Properties 12 A. Zecchina, S. Coluccia, L. Cerruti and E. Borello, J. Phys. Chem., 1971, 75, 2783. 13 A. Zecchina, S. Coluccia, E. Guglielminotti and G. Ghiotti, J. Phys. Chem., 1971, 75, 2790. 14 I. Kostov, Mineralogy (Nauka i Izkustvo, Sofia, 1973). 15 J. C. Lavalley, J. Saussey and T. Rais, Proc. VIth Soviet-French Seminar on Catalysis, Moscow, 1983, 16 K. Hadjiivanov and A. Davydov, Kinet. Kafal., in press. 17 P. Vergnon, J. M. Hermann and S . J. Teicher, Russ. J. Phys. Chem., 1978, 52, 3021. 18 V. Lorenzelli and G. Busca, Muter. Chem. Phys., 1985, 34, 261. 19 B. F. Ormont, Structure oflnorganic Compounds (Gosudarstveno Izdatelstvo NT literaturoi, Moscow- 20 K. Hadjiivanov, D. Klissurski and A. Davydov, Ann. Univ. Sofa Fac. Chim., 1985, 79, in press. p. 97. Leningrad, 1950). Paper 611738; Received 27th August, 1986

ISSN:0300-9599    

DOI:10.1039/F19888400037

出版商:RSC    

年代:1988

数据来源: RSC

摘要:

J. Chem. SOC., Faraday Trans. I , 1988, 84(1), 41-56 A Polarimetric and llB and 13C Nuclear Magnetic Resonance Study of the Reaction of the Tetrahydroxyborate Ion with Polyols and Carbohydrates J. Graham Dawber," Stuart I. E. Green and (in part) John C. Dawber and Sundus Gabrail Department of Chemistry and Biology, North Staffordshire Polytechnic, Stoke-on-Trent ST4 2DE The reaction of the tetrahydroxyborate ion, B(OH);, with 31 polyols and carbohydrates in aqueous solution has been studied by complementary investigations involving polarimetry and llB and 13C n.m.r. spectroscopy. The calculation of equilibrium constants for the complexation of the carbohydrates with B(0H); from the polarimetry results, using a previously derived equation, was only partially successful owing to the presence of multiple equilibria in a number of cases.The presence of several complexed species in these reactions is demonstrated by llB n.m.r. studies. A calibration method devised to relate peak area in llB n.m.r. spectra with concentration has been used to calculate equilibrium constants for the various equilibria present, and an attempt has been made to rationalise the equilibrium constants with the molecular structures of the substrates. In the majority of cases the 13C n.m.r. spectra confirm the polarimetric and llB n.m.r. studies, and in most cases allow a more specific identification of the reaction sites in the polyols/carbohydrates to be made. The use of polyhydroxy compounds as a means of increasing the strength of boric acid for its titration has been known for many years and the reaction has also been used to characterise carbohydrates.However, the stoichiometry of the complexes formed has often been uncertain.1-6 It is known that the reaction with polyols is much more pronounced when the tetrahydroxyborate ion, B(OH),, is used instead of boric acid itself.'* While the complexation of borate across adjacent hydroxy groups in the polyol is commonly assumed, it is now evident that complexation across alternate hydroxy groups is also possible, as demonstrated by 'H n.m.r. spectros~opy,~ "B n.m.r. spectroscopy, lo circular dichroismll and polarimetry . l1 At high ratios of [polyol]/ [borate] it is possible to observe two molecules of polyol to one molecule of boratelO in the complex. In the case of cyclic myo-(meso) inositol it is thought that three hydroxy groups on alternate carbon atoms can become involved in a tridentate complex with the B(0H); ion.12 While our previous ''B n.m.r.work was in progress, unknown to us, Kieboom and coworker^^^^ l4 were carrying out similar studies, and the two pieces of work agree very well, although their results were published while ours were still in press." The work of Kieboom and c o ~ o r k e r s ~ ~ ~ l4 has since taken a different direction from ours in that they have carried out an in-depth study of polyol carboxylate salts and their reactions with borate. Our work, on the other hand, has included a study of a wider range of non-ionic polyols and also a range of monosaccharides and two disaccharides, since our interest was in the reaction of borate with un-ionised species.The techniques used have been polarimetry, "B n.m.r. and l3C n.m.r., and these results are reported here. We have modified our method of calculating the equilibrium constants of complexation from the 4142 Polarimetry and "B and 13C N.M.R. "B n.m.r. studies and the recalculated values of our previous worklo are also presented. Experimental Materials The materials used were all of GPR quality (except glucose and sucrose which were AR), and were used without further purification. The materials were as follows (structures are given in table 1) : ethane- 1,2-diol (l), propane- 1,2-diol(2), propane- 1,3-diol(3), glycerol (4), butane- 1,2,4-triol (5), meso-erythritol (6), threitol (7), xylitol (8), D-arabitol (9), adonitol (ribitol) (lo), dulcitol (galacticol) (1 l), mannitol (12), sorbitol (13), cis- cyclohexane- 1,2-diol (14), trans-cyclohexane- 1,2-diol (15), cyclohexane- 1,3-diol (16), cyclohexane-r- 1 ,c-3,c-5-triol (17), myo(rneso)-inositol (18), dihydroxydioxane (19), sodium ascorbate (20), a-methyl glucoside (21), D( + )-glucose (22), D( +)-galactose (23), D( +)-mannose (U), L( +)-rhamnose (25), D( +)-xylose (26), D( -)-arabinose (27), L( +)-arabinose (28), fructose (29), sucrose (30) and maltose (31).The NaB(OH), solution (stock solution, 2.5 mol dm-3) was prepared by direct reaction of A.R. NaOH and H,B03, the pH was then adjusted to 12.5, and this solution was diluted as required. Preparation of Borate of Cyclic Polyols (a) cis-Cyclohexane- 1,2-diol Borate Equimolar quantities of the diol and NaB(OH), were mixed in solution and heated to ca.70 "C. Ethanol was then added carefully until the solution was just cloudy. The solution was next heated until it was clear again and then allowed to cool very slowly and left to stand for at least 24 h. White, well formed shiny leaflets of the borate crystallised out (decomp. above 300 "C after loss of water at lower temperatures; analysis gave C, 29.8 YO ; H, 7.6 YO ; the trihydrate C,H,,O,BNa requires C, 30.5 YO ; H, 7.7 %). (b) Cyclohexane-r-1 p3,c-j- triol Borate The borate of this triol was prepared as for (a). Well formed lustrous flakes were obtained (m.p. > 340 " c ; analysis gave c , 30.8 YO ; H, 6.9 % ; the trihydrate C,H,,O,BNa requires C, 30.8 YO ; H, 6.8 %).(c) myo-Inositol Borate This was prepared as for (a), but on cooling, the complex initially separated from the solution as an oil, which slowly solidified to a crystalline solid. This product appears to retain more water of crystallisation than the rather better formed crystals in (a) and (b), but of course contains many more hydrophilic centres (decomp. from 320 "C after losing water at much lower temperatures; analysis gave C, 20.5 % ; H, 6.7 YO ; the heptahydrate C,H,,O,,BNa requires C, 20.4 YO ; H, 6.8 YO). Polarimetric Measurements with the Carbohydrates Solutions for the polarimetric measurements were made up containing 0.25 mol dm-3 of carbohydrate and the concentration of NaB(OH), varied from 0 to 1.25 mol dm-3. The optical rotations of the solutions were measured with a Bellingham and Stanley model A photoelectric polarimeter using a 100 mm pathlength tube.Angular rotations couldJ . G. Dawber, S. I. E. Green, J . C. Dawber and S. Gabrail 43 Table 1. Structures of compounds studied HO HO HO WOH HO Ho* * HO OH HO HO OH OH OH HO (9) Horn &OH HO HO HO OH HO OH OH OH (14) OH (17) OH (15) OH 6.;.;, OHHo 0 Na OH (19) \ OH44 Polarimetry and "I3 and I3C N.M.R. Table 1. Structures of compounds studied (con?.) OCH3 (21) CH,OH HovTj-ti - - bo (OH,CH,OH ) OH (OH , CHZOH) OH (29) a Labelling system for carbohydrate carbon atoms. be estimated to 0.002". The measurements were carried out at a wavelength of 435.8 nm using a low-pressure Hg lamp with the unwanted wavelengths filtered out by a solution of NaNO, and a cobalt blue filter.15 This wavelength was used rather than the sodium D-line since it provided larger differences in optical rotation between successive solutions, thereby giving better discrimination. The temperature was 20 1 "C.J.G. Dawber, S. I. E. Green, J. C. Dawber and S. Gabrail 45 N.M.R. Measurements (a) "B N.M.R. The llB n.m.r. spectra were measured using a Jeol FX-90Q Fourier-transform spectrometer (lH resonance at 89.55 MHz, llB resonance at 28.75 MHz) using a tip angle of 45" and a pulse repetition time of 2 s. The spectral responses over 4500 Hz were acquired into 8 K data points and zero-filled to 16 K data points and an experimental broadening of 0.3 Hz applied prior to Fourier transformation. All solutions were made up in water with the spectrometer locked onto a D20 capillary.Referencing of the "B chemical shifts was made relative to H,BO, in water as 6 = 0.0 ppm. The concentration of NaB(OH), used was 0.25 mol dm-, and the concentrations of polyol/carbohydrate ranged from 0 to 1.0 mol dm-3. When the B(OH), ion complexes with polyols the llB n.m.r. signal appears at characteristic positions for the various types of 11* 1 3 7 l4 The formation of complexes involving vicinal hydroxy groups is accompanied by downfield shifts (for 1 : 1 and 1 : 2 complexes), and accompanied by upfield shifts for complexes involving hydroxy groups on alternate carbon atoms. The calculation of equilibrium constants of formation for the various complexes has previously been based upon the integration of the areas of the various peaks in the llB n.m.r.spectra. Implicit assumptions in this approach are (i) that the area of a given llB peak is directly proportional to the concentration of that boron-containing species, and (ii) that the same relationship holds between peak area and concentration for all the different boron species. The half-widths (A$ of the various llB peaks in this study ranged between 4 and 60 Hz, and, if one approximates the effective quadrupolar relaxation time as (zAv$',l6 the relaxation times of the various boron species should lie within the pulse repetition time of 2 s. Nevertheless, it was decided to check the linearity of the peak area with concentration for the B(0H)i species in two series of experiments in which the n.m.r. spectra of various concentrations of NaB(OH), were measured simultaneously with a fixed concentration of boron trifluoride etherate in a capillary in the n.m.r.tube. In one series neat boron trifluoride etherate was used, whereas in the other series a solution of the etherate in CHC1, (ca. 2 mol drn-,) was placed in the capillary. The area of the B(OH), ion peak (the borate solution containing ca. 10% D20 was placed in the main n.m.r. tube) was measured relative to that of the boron trifluoride etherate. This was done for a wide range of borate concentrations. Fig. 1 shows plots of the ratio of the areas of the two peaks as a function of B(OH), concentration (the etherate being effectively an analytical internal standard) and it can be seen that there is good linearity, indicating that the peak area is proportional to boron concentration.There is still, however, the question of the relationship between the different peaks for the various borate species, namely, B(OH),, the 1 : 1 complexes (BP-) and the 1 :2 complexes (Bpi). This feature is important since in the normal experimental runs the proportion of each boron-containing species is taken as a fraction of the total boron concentration as calculated from its percentage of the total area of the llB n.m.r. signals. A further two series of experiments were performed using the boron trifluoride etherate capillary as the fixed concentration standard. In one set of experiments various proportions of sorbitol were added to 0.25 mol dm-, NaB(OH), in order to compare the relative sizes of the peaks from B(OH),, BP- and BP, (P = polyol).In a second set of experiments xylitol was used as the polyol. For the sorbitol experiment it was found that the relative peak area sizes for the same boron concentration for the species BP,, BP- and B(0H); were in the ratio 1.00:0.79:0.68, and for the corresponding xylitol experiment the values were 1 .OO : 0.76 : 0.62, which was considered to be in reasonable agreement. The data from the sorbitol experiment were chosen for the purpose of correcting the relative peak areas for the calculations of the equilibrium constants of the various equilibria (see later).46 Polarimetry and "B and 13C N.M.R. [ NaB(OH)4]/mol dm-3 Fig. 1. B(0H); calibration: 0, neat etherate standard; x , diluted etherate standard. (b) 13C N.M.R.The proton-decoupled 13C n.m.r. spectra were measured on a Jeol FX90Q spectrometer (13C resonance at 22.49 MHz) using a tip angle of 30" and a pulse repetition time of 1 s. The spectral responses over 4500 Hz were acquired in 8 K and zero-filled to 16 K data points and an experimental broadening of 0.7 Hz was applied prior to Fourier- transformation. The solutions were all made up in H,O and the instrument was locked onto a D,O capillary in the n.m.r. tube. Referencing was relative to TSP at 6 = 0.0 ppm. The spectrum of each polyol or carbohydrate was measured at 1 mol dm-3 in H20 and 1 mol dmd3 in 2.5 mol dm-3 NaB(OH), solution. Results and Discussion Polarimetry The optical rotation results were converted to molar optical rotation, [Y], by E\y] = a/d, where a is the optical rotation in a polarimeter tube of 1 metres for a carbohydrate concentration of c mol m-3, giving units of O m2 mol-1 for [Y].The values of [Y] at 436 nm for the various carbohydrates in water are given in table 2. For each carbohydrate, when borate was added, the change in molar rotation, A["], from the value in water was calculated and these values are plotted as a function of NaB(OH), concentration in fig. 2-4. It can be seen that in all cases the reaction of the chiral carbohydrates with the B(OH), ion was evident from the changes in molar rotation. In most cases the variation of A w l changes simply with NaB(OH), concentration, but for L( +)-rhamnose and D( +)-xylose the behaviour can be seen to be more complicated In previous work8*l1 a method was devised for calculating the equilibrium constant, K,, for the formation of complexes of polyols using polarimetry results.Changes in the molar optical rotation are dependent upon the extent of complexation, the stoichiometry (fig- 4).J . G. Dawber, S. I. E. Green, J . C. Dawber and S. Gabrail 41 Table 2. Optical rotation results anomeric composition pyranose (YO) a /3 furanose (YO) ["I/' m2 molt1 K,/dm3 mol-I a-methyl glucoside D( + )-glucose D( + )-galactose D( + )-mannose L( + )-rhamnose D( + )-xylose D( - )-arabinose L( + )-arabinose fructose sucrose maltose sodium ascorbate 100 36 27 67 34 33 63 63 3 37 - - 0 64 73 33 66 67 34 34 90 67 - - 0.5954 0.1830 0.2753 0.05 12 0.0273 0.0567 0.2993 0.4379 0.90 18 0.439 1 -0.3046 - 0.3208 1.3 18.9, 2.9 a a a a a a a 0.7, 0.5 6.1, 0.9 4.7 a Unable to be evaluated from the polarimetric results with any great certainty, the results giving curves rather than the required straight lines.1 I I I I 1 2 3 4 (borate]/[ sugar] Fig. 2. Changes in molar optical rotation with added NaB(OH), + , a-methyl glucoside; x , D( +)- glucose ; A, D( + )-mannose ; 0, D( + )-galactose.48 Polarimetry and "B and 13C N.M.R. 300 200 100 4 I I i E O 2 e m ;= Q -100 -200 - 300 X I I I I 1 2 3 4 [borate]/ [ sugar] Fig. 3. Changes in molar optical rotation with added NaB(OH),: 0, D( -)-arabinose; x , L( +)- arabinose ; A, fructose; 0, sodium ascorbate. of the complex, and the molar rotation of the complex (or complexes) compared to that of the original polyol/carbohydrate. The equations derived assumed the formation of a 1 : 1 complex and were successfully applied to systems involving a number of anions with sorbitol and mannitol* and also a chiral diol.ll However, in the case of B(0H); and sorbitol and mannitol the results did suggest the possibility of more than one complexed species being present.In the determination of K, for the carbohydrates of the present work, by far the simplest behaviour was found for a-methyl glucoside. For the remainder of the carbohydrates the evaluation of K, proved to be less simple. Instead of a simple linear plot being obtained from the method,8 for some carbohydrates the graphs consisted of two linear portions leading to two values for K,. For a number of other cases (galactose, mannose, arabinose, xylose, rhamnose and fructose) K, could not be evaluated satisfactorily, and the reasons for this are discussed later.The K, values for the carbohydrates for which the above method of evaluation was successful are presented in table 2. Most of the carbohydrates studied in this work exist in solution as a mixture of anomers, e.g. the a- and B-pyranose, and the a- and 8-furanose structures. The initial polarimetry studies showed that whatever the anomeric form of the initial solid carbohydrate, the equilibrium anomeric mixture in solution was achieved very rapidly (< 30 s) on addition of a small amount of NaB(OH), solution. The proportions of theJ . G. Dawber, S. I. E. Green, J. C. Dawber and S. Gabrail 49 I I I I I 1 2 3 I [borate I/[ sugar] Fig. 4. Changes in molar optical rotation with added NaB(OH), 0, D(+)-xylose; A, L(+)- rhamnose; x , sucrose; 0, maltose.a- and 8-pyranose structures make up the majority of the anomeric mixture, with the furanoside structures accounting for only 1-2 % of the total composition1' (table 2). Hence we assume that complexation with the B(0H); ion involves principally the a- and 8-pyranoside structures. The structural difference between glucose itself and that of a-methyl glucoside is the availability for complexation of the OH group on the C , carbon atom in glucose along with its adjacent OH group on C,, hence one would expect a different and enhanced value of K, and this was found to be the case. A similar relationship exists for the disaccharides sucrose and maltose [ i e . a-glucopyranosyl fructose and 4-0-a-(D)- glucopyranosyl D-glucose], with the maltose complexing considerably more with B(OH), than with sucrose, a significant difference between the structures of the two sugars being the a- and 8-OH group on the C , carbon atom of maltose.In situations where complex formation can occur at many sites, including adjacent and alternate carbon atoms (see later), it is not surprising that the optical rotation results in some of the cases do not allow an evaluation of K, from a method based upon a simple model of complexation. * "B N.M.R. For the purposes of comparing the complexation of the B(0H); ion with the various polyols/carbohydrates we have considered three general reaction equilibria. Two of these represent complexation across adjacent carbon atoms and the other represents complexation across alternate carbon atoms (including the case of tridentate coordination of the B atom): K , B- + P f Bpidj.K2 BP,,. + P f B P i ~ j . K3 B- + P f Bpil,.50 Polarimetry and llB and 13C N.M.R. [B- represents B(OH),, P represents polyol/carbohydrate, and adj. and alt. represent reaction across adjacent and alternate carbon atoms, respectively]. With higher-field llB n.m.r. l4 it is possible to observe a more comprehensive range of equilibria involving various conformers, but these were only just discernible in our studies at lower field. In calculating the various equilibrium constants we used the weighting of the peak areas found in the calibration procedure with sorbitol. If the fractional areas of the total peak area fyr "B are a, p, y and 6 for the species B-, Bpidj., Bpiadj.and Bpil,., respectivelyJand the weighted peak area factors from the calibration study are a, b, c and d (d was assumed to be the same as b), and if m, is the total concentration of borate in all its forms and rn, is the total polyol concentration in all its forms, then and [PI = m,-(bp+2cy+d6)ml/D where D = (aa+bp+cy+dd). From the sorbitol calibration experiment a = 0.68, b = 0.79, c = 1.00 and d = 0.79, and hence the equilibrium constants K,, K, and K, can be calculated for the various polyol/carbohydrate - borate systems from the relative integrated peak areas (a, p, y, 6 ) in the "B n.m.r. spectra. For the purposes of comparison the values of K,, Kz and K3 were calculated at a total polyol concentration (m,) of 0.5 mol dm-3 and a total borate concentration (m,) of 0.25 mol dm-3, and the values obtained are presented in table 3.In the small area of duplication of other the trends in our values of the equilibrium constants are similar. For the systems where Kc could be evaluated from the polarimetry results the corresponding values of the equilibrium constants from the llB n.m.r. results agree at the same level of compatibility as previously found.lOvll It can be seen from table 3 that for the various linear polyols the magnitude of the values of the equilibrium constants are related to the number of OH groups available for complexation. In general the magnitudes are in the order K, > K, > K3, showing that complexation across adjacent OH groups is easier than that across alternate OH groups.We are of the opinion that for the linear polyols there is relatively ease of rotation of the CH,OH groups at the end of the chains, thus accounting for the low K, values for the lower members of the series. When one reaches the C, members, meso-erythritol and threitol (and to some extent the trio1 glycerol), there is a considerable increase in the value of K,, and it is for these compounds that there is relatively more restriction to movement of the OH groups in the 'inner' CHOH groups, and it is these which are likely to be involved in the complexation reaction with the B(OH), ion. For the C, polyols, xylitol and D-arabitol, there is a further increase in K,. In the case of xylitol there is some competition for complexation across alternate OH groups (K3 = 17) and inspection of its structure (table 1) shows that the OH groups on C, and C , are favourably situated for complexation. This is not the case for D-arabitol, which has no value for K3 but a larger value of K l .In the case of adonitol, K, (at 32) is the lowest for the C, polyols, but K3 is considerable (20). This enhanced value for K3 may be due to the fact that adonitol has a number of alternate OH groups favourably disposed with respect to each other, namely Cl:C3, C,:C4 and C3:C,. In these three polyols it is thought that the most favourable positions for complexation involve adjacent OH groups which are disposed to each other at an angle of 60". Of course the conformations shown in table 1 are not fixed and rotation about single C-C bonds willJ.G. Dawber, S. I. E. Green, J. C. Dawber and S. Gabrail Table 3. llB N.m.r. results for determination of equilibrium constants of complexation (dm3 mol-l) 51 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 ethane- 1,2-diol propane- 1,2-diol propane- 1,3-diol glycerol butane- 1,2,4-triol meso-erythri to1 threi to1 xylitol D-arabitol adonitol dulcitol mannitol so rbi to1 cyclohexane- 1,2-diol trans-cyclohexane- 1,2-diol cyclohexane- 1,3-diol cyclohexane-r- 1 ,c-3,c-5-triol myo-inositol 2,3-dihydroxydioxane sodium ascorbate a-methyl glucoside D( + )-glucose D( + )-galactose D( + )-mannose L( + )-rhamnose D( + )-xylose D( - )-arabinose L( + )-arabinose fructose sucrose maltose 3 .O 4.7 - 37 70 78 126 540 32 1000 232 540 6.2 1.4 - - - 4.4 3.7 2.3 525 17 113 I99 44 347 246 246 235 - 12.3 0.4 0.8 3.9 1.7 - 12 12 15 68 5 32 68 - 1.8 - - - 1.9 0.4 8 I5 6 3 31 32 26 111 58 - - 2.2 occur, but the structures are likely to represent the most energetically favoured conformations.For the C, linear polyols the magnitude of Kl is in the order mannitol < sorbitol < dulcitol. In the case of mannitol there is a favourable pair of OH groups on C, : C , ; in the case of sorbitol there are two pairs of potential OH groups (but one OH group is common), i.e. C, : C , and C,: C,; whereas in the case of dulcitol there are two separate pairs of OH groups, i.e. C, : C , and C, : C, (there was some evidence from the 13C n.m.r. data that C,:C, OH groups might also be involved).Hence the values of Kl are approximately compatible with the molecular structures. cis-Cyclohexane- 1,2-diol was found to form a complex with B(OH),, but for the trans isomer we found no evidence of reaction. The trans-1,2-diol in the neat liquid can form intramolecular hydrogen bonds between the adjacent OH groups when both are in equatorial positions. However, we feel that in aqueous solution the opposite chair configuration could be favoured, since in this conformation the two OH groups become axial, thus allowing more access for solvation by water molecules. In this conformation complexation of the trans isomer with B(OH), would be impossible. The cyclohexane-1,3-diol (mixture of cis and trans) showed no evidence of complexation with borate.By contrast, however, cyclohexane-r- 1 ,c-3,c-5-triol showed52 Polarimetry and "B and 13C N.M.R. (a 1 -14.1 -17.4 -18.5 -17.4 -13.6 -19.2 Fig. 5. llB n.m.r. spectra of polyol borate complexes in water: (a) cis-cyclohexane-l,2-diol, (b) cyclohexane-r- 1 ,c-3,c-5-triol, (c) myo-inositol. considerable reaction with B(0H); (K3 = 39), but in this case the reaction is thought to involve a tridentate complex with the b0r0n.l~ In the case of myo-inositol the largest value of equilibrium constant is for K3, and again this is likely to form a tridentate complex12 similar to cyclohexane-r-1 ,c-3,c-5-triol. Nevertheless, our "B n.m.r. results showed that some reaction also occurred across adjacent OH groups of the inositol. The complexes of cis-cyclohexane- 1,2-diol, cyclohexane-r- 1 ,c-3,c-5-triol and myo-inositol were isolated (see Experimental) and the llB n.m.r.spectra of the compounds dissolved in water are shown in fig. 5. Here it can be seen that the complexes are extensively dissociated into polyol and B(0H); (6 = - 17.4 ppm). In the case of myo-inositol the small amount of bidentate complex involving adjacent OH groups is also evident at a peak position of 6 = - 13.6 ppm. The extent of complexation of 2,3-dihydroxydioxane with B(OH), is very extensive compared to that of cis-cyclohexane- 1,2-diol. It is possible that the two OH groups are much more acidic in the dioxane compound, owing to the inductive effects of the ring oxygen atoms rather than any stereochemical/configurational influence. Sodium ascorbate, by contrast, is rather more like cyclohexane-1,2-diol in its extent of complexation with B(OH),, The K values for the carbohydrates (compounds 21-31) are shown in table 3 and it can be seen that complexation with borate is extensive in many cases.For four of the carbohydrates (glucose, xylose, mannose and maltose) the manner of complexation is comprehensive, involving the equilibria for K,, K, and K3. In the particular case of xylose its value of K3 is considerable (in addition to K , and K,) and this coincides with the complicated optical rotation behaviour observed on addition of borate. The values of K for the pairs methyl glucoside and glucose, and sucrose and maltose have a similar relationship, and this similarity was also observed in the optical rotation measurements.As discussed earlier, this relationship can be rationalised with theJ. G. Dawber, S. I. E. Green, J. C. Dawber and S. Gabrail 53 availability of an OH group on the C, carbon atom, giving rise to enhanced complexation. For mannose and glucose the configurations about C,, C,, C,, C, are the same, but that for galactose is different. It may be significant that for mannose and glucose the reaction with B(OH), can also involve alternate OH groups to a small extent, whereas in the case of galactose the reaction involves only the equilibria for Kl and K2. The cyclic pyranose structures of the pentoses xylose and the two isomeric arabinoses have similarities to the pyranoside form of fructose, and it is for these four compounds that the highest K values were observed amongst the carbohydrates as a group.Unfortunately there does not seem to be a prima facie case of a simple correlation between the individual molecular structures and their corresponding values of the equilibrium constants, K. 13C N.M.R. The results of the 13C n.m.r. spectra are given in table 4 [the complete table, including chemical shift assignments, has been deposited as supplementary publication no. SUP 56699 ( 5 pp)].t The spectral assignments of the compounds were made by utilising information from several sources. Our chemical shift values are slightly different from those in the literature since they are all relative to TSP at 6 = 0.0 ppm, but all the assignments are consistent with the other published data for the uncomplexed polyols/ carbohydrates. The 13C spectra of the compounds dissolved in 2.5 mol dm-, NaB(OH), (a borate : carbohydrate ratio of 2.5) invariably showed broadening of all the resonances, with the carbon atoms involved in the complexation being affected most. Detailed comments relating to the spectra and their interpretation are given in table 4.For the linear polyols (compounds 1-13) the earlier members (ethane diol, propane- 1,2-diol, propane- 1,3-diol, glycerol, butane- 1,2,4-triol) all show evidence of the terminal CH20H group(s) being involved in complexation. However, when the chain length is more extended (compounds 6-13), complexation with B(0H); favours the inner CHOH groups, which are likely to have less freedom of rotation than the terminal CH20H groups. This would account for the large increase in complexation as the chain-length of the polyol is increased (table 3).Adonitol was found by the llB n.m.r. studies to be involved considerably in complexation with OH groups on alternate C atoms, in addition to the reaction across adjacent carbon atoms, and this extra feature shows itself in the very complicated 13C n.m.r. spectrum of this compound in 2.5 rnoldm-, NaB(OH),. The cyclic polyols cis-cyclohexane- 1,2-diol and cyclohexane-r-1 ,c-3,c-5-triol have relatively simple 13C n.m.r. spectra in water and also as complexes with B(0H);. The results for the diol show simple complexation, while those for the trio1 show unequivocal evidence for the formation of a tridentate complex involving all three OH groups. For myo-inositol, however, the 13C n.m.r. spectrum in borate showed evidence of bidentate as well as tridentate complexation, in agreement with the "B results.The 13C n.m.r. data for the carbohydrates (21-31) are also given in table 4. As might be expected, the simplest behaviour was observed for a-methyl glucoside, where the C, carbon atom has no available OH group and the conformation is also simplified by the absence of other anomers. The spectra for the other carbohydrates confirmed our suspicions that several of the available OH groups can compete for complexation with the borate, and the particular findings are summarised in table 4. In a number of cases the spectra of the carbohydrates in NaB(OH), were so complex that the individual resonances could not be assigned. This strongly suggests in these cases that a variety of complexed species were present together in solution and these cases invariably were t See Notice to Authors, J.Chem. SOC., Faraday Trans. 1, 1988, 84, January issue.54 Polarimetry and and I3C N.M.R. Table 4. Complexation trends inferred from n.m.r. data compound comments on complexation with B(0H); 1 ethane- 1,2-diol 2 propane- 1,2-diol 3 propane- 1,3-diol 4 glycerol 5 butane- 1,2,4-triol 6 meso-erythritol 7 threitol 8 xylitol 9 D-arabitol 10 adonitol 11 dulcitol 12 mannitol 13 sorbitol 14 cis-cyclohexane- 1,2-diol 17 cyclohexane-r- 1 ,c-3,c-5-triol 18 myo-inositol 19 2,3-dihydroxydioxan 21 methyl glucoside 22 glucose Simple complexation across C, : C,. Simple complexation across C, : C,. Complexation across C, : C,.Upfield shift of C, resonance. From the 'lB n.m.r. results the greatest complexation involves adjacent C atoms. The three 13C signals indicate C, : C, and C, : C,. From the llB n.m.r. results the greatest complexation involves adjacent C atoms. The 13C indicates C,:C,. Complexation mainly across C,:C,. If C,:C, or C,:C, involved then there would be 4 main peaks. Complexation mainly across C,:C,. If C,:C, or C,:C, involved then there would be 4 main peaks. The inner CHOH groups affected most. Complexation across C, : C, or C, : C,. The C, and C, signals are slightly different. Complexation across C,:C, and C,:C, would produce more lines than for xylitol. The llB n.m.r. suggests considerable reaction across alternate C atoms as well as adjacent positions. If C,:C,, C, : C,, C, : C,, C, : C, and C, : C, positions were involved then this would give possibility of 18 resonances. Principally 3 lines convert into 3 lines indicating C,:C, complexation.The weaker shoulders, however, suggest some C, : C, and C, : C, involvement. Reaction mainly across C, : C,, these two OH groups have the closest disposition. Reaction across C, : C,, C, : C,, C, : C,. Some suggestion that C, may be involved also. Simple complexation across C, : C,, with some different conformations about C, and C,. Formation of the complex produces just two lines and thus must involve a tridentate complex, any other complex would produce more than two resonances. The central 13C signal goes upfield, cJ: propane- 1,3-diol. In excess borate there is probably considerable bidentate complex present in addition to the majority tridentate.The prepared complex when dissolved in water gave (from the llB n.m.r.) 50.3 YO tridentate, 1 1.4 YO bidentate (adjacent OH groups, C,:C, or C,:C,) and 28.3% uncomplexed polyol. Such a mixture should give rise to 14 lines and the 13C n.m.r. spectrum of the prepared complex in water did in fact show 14 signals, at 77.9, 77.7. 77.3, 76.2, 75.4, 75.2, 74.1, 73.7, 71.7, 70.6, 69.0, 66.7, 66.4, 65.3, but of course many of these were overlapping and this influenced their positions. Reaction at C,:C,, although additional small lines at 107.1, 101.9, 99.6, 95.2, 92.7 and at 71.5, 64.6, 63.9, 63.5 suggest that several conformations may be involved. Intensities of C, and C, affected most :. adjacent complexation across C, : C,.Possibly C, has axial conformation. Spectrum suggests complexation mainly across C, : C, and C,:C, and possibly C,:C,. Possibly C, has axial conformation, and molecule in &conformation.J . G. Dawber, S . I. E. Green, J. C. Dawber and S. Gabrail 55 Table 4. Complexation trends inferred from n.m.r. data (cont.) _ _ compound comments on complexation with B(0H); 23 galactose 24 mannose 25 rhamnose 26 xylose 27, 28 arabinose 29 fructose 30 sucrose 31 maltose Spectrum suggests complexation across C, : C, mainly with molecule in a-conformation. In borate all the lines are affected and all broadened. Although the complexation will mainly be across adjacent C atoms the IIB n.m.r. indicated considerable alternate C-atom involvement.Lines at 105.2, 100. I , 97.4, 96.8, 85.4, 84.1, 78.6, 77.2, 73.6, 72.8, 70.7, 67.2, 62.2. Polarimetry showed complex behaviour. 21 new lines produced, many of which occur as pairs: 187.9, 105.4, (99.5, 99.0), (97.1, 96.6), 95.9, (85.1, 84.7), 81.3, (79.3, 78.3, 77.3), 69.5, 68.6, 34.5, (22.7, 22.4), (21.3, 20.9), 15.4. Thus there must be a variety of complexes present, including the possibility of open-chain material (n.b. the line at 187.9 which is likely to be the C=O group). The CH, signal at high field appears at several positions, suggesting a number of different species. Although the llB n.m.r. indicates principally complexation across adjacent C atoms, the optical rotation behaviour was unusual. Spectrum suggests C,, C,, C,, C, involved. The 'lB n.m.r.suggests complexation across alternate as well as adjacent C atoms. Polarimetry gave unusual behaviour. Spectrum suggests C,, C,, C, (C, possibly) involved. The llB n.m.r. indicates predominantly adjacent C atoms involved. Polarimetry could not give value of K,. Complexation probably across C, : C, and C, : C,, mainly in the pyranoside p-form. Main interaction, which is slight, appears to involve C, and C,. The IlB n.m.r. shows little interaction and com- plexation involves alternate C atoms. llB n.m.r. suggests reaction across adjacent and alternate C atoms (more than in sucrose). The 13C n.m.r. spectrum suggests reaction at C,, C,, Ci, Ch, but it is difficult to decide which of the adjacent C atoms are involved. The line at 188.8 ppm suggests that some of the material may be in the open-chain configuration.those for which a value of K, could not be evaluated from the polarimetric results and which also exhibited complex optical rotation behaviour. Of the two disaccharides studied, the 13C n.m.r. data for sucrose were the easier to interpret and suggested (as with the "B n.m.r. data) that the low extent of interaction with B(OH), ion involves complexation across alternate carbon atoms. The spectrum of maltose in NaB(OH), was less easy to interpret, except that it was possible to observe alternate OH-group involvement in addition to that of adjacent groups, and the presence of an n.m.r. line in the C=O region suggests that there may be a small amount of complexation involving the open-chain material. Conclusions Complementary studies of polarimetry, "B n.m.r.and 13C n.m.r. show that the B(OH), ion has almost universal ease of complexation with polyols and carbohydrates. The results show conclusively that in many cases the reaction involves various56 Polarimetry and 'I% and 13C N.M.R. competing equilibria in which several of the OH groups in the polyol/carbohydrate may participate, and that the interpretation of the behaviour is greatly facilitated by llB n.m.r. spectroscopy. The previously developed method of evaluating equilibrium constants from polarimetric results was of limited use for the complex equilibria in these systems. We thank Mr P. Doughty of the Mining Engineering Department, N.S.P., for carrying out the C and H analyses, and the S.E.R.C. for funds towards the cost of the polarimeter.References 1 J. Boiseken, Adv. Carbohydr. Chem., 1949,4,189; 1949,12,81; J. P. Sickels and H. P. Schultz, J. Chem. 2 R. F. Nickerson, J. Inorg. Nucl. Chem., 1968, 30, 1447; 1970, 32, 1400. 3 G. W. Campbell, J. Inorg. Nucl. Chem., 1969, 31, 2626. 4 H. B. Davis and C. J. B. Mott, J. Chem. Soc., Faraday Trans. 1, 1980, 76, 1991. 5 R. Larsson and G. Nunziata, Acta Chem. Scand., 1970, 24, 2145. 6 M. Mazurek and A. S. Perlin, Can. J. Chem., 1963, 41, 2403. 7 J. G. Dawber and D. H. Matusin, J. Chem. Soc., Faraday Trans. I , 1982, 78, 2521. 8 J. G. Dawber and G. E. Hardy, J. Chem. Soc., Faraday Trans. 1, 1984, 80, 2467. 9 R. E. Moore, J. J. Barchi and G. Bartolini, J. Org. Chem., 1985, 50, 374. 10 J. G. Dawber and S. I. E. Green, J. Chem. Soc., Faraday Trans. I , 1986, 82, 3407. 11 J. G. Dawber, J. Chem. SOC., Faraday Trans. 1 , 1987, 83, 771. 12 S. J. Angyal and D. J. McHugh, J. Chem. Soc., 1957, 1423; S . J. Angyal, J. E. Klavins and J. A. Mills, Aust. J. Chem., 1974,27, 1075; P. J. Garegg and K. L. Lindstrom, Acta Chem. Scand., 1971,25, 1559; R. M. Williams and R. H. Attala, in Solution Properties of Polysaccharides, ed. D. A. Brant (American Chemical Society, Washington D.C., 1981). 13 M. Van Duin, J. A. Peters, A. P. G. Kieboom and H. Van Bekkum, Tetrahedron, 1984, 40, 2901; M. Makkee, A. P. G. Kieboom and H. Van Bekkum, Reel. Trav. Chim. Pays-Bas, 1985, 104, 230; M. Van Duin, J. A. Peters, A. P. G. Kieboom and H. Van Bekkum, Tetrahedron, 1985, 41, 3421 ; M. Van Duin, J. A. Peters, A. P. G. Kieboom and H. Van Bekkum, Reel. Trav. Chim. Pays-Bas, 1986, 105, 1986. Educ., 1964, 41, 343; J. Boiseken and N. Vermaes, J. Phys. Chem., 1931, 35, 1477. 14 M. Van Duin, Doctoral Thesis (Technological University of Delft, Delft University Press, 1986). 15 J. G. Dawber, J. Chem. SOC., Faraday Trans. 1, 1978, 74,960. 16 R. K. Harris, Nuclear Magnetic Resonance Spectroscopy (Pitman, London, 1983), p. 71. 17 R. J. Ferrier and P. M. Collins, in Monosaccharide Chemistry, Penguin Library of Physical Sciences: Chemistry (Penguin Books, Harmondsworth, 1972) ; J. F. Stoddart, in Stereochemistry of Carbo- hydrates (Wiley, New York, 1971). 18 W. Voelter, E. Breitmaier, G. Jung, T. Keller and D. Hiss, Angew. Chem. Int. Ed. Engl., 1970, 9, 803; K. Bock and H. Thogersen, Ann. Rep. NMR Spectrosc., 1982, 13, 1; M. Voelter, V. Bilik and E. Breitmaier, Collect. Czech. Chem. Commun., 1973,38, 2054; T. A. W. Koerner, R. J. Voll, L. W. Cary and E. S. Younathan, Biochem. Biophys. Res. Commun., 1978, 82, 1273; L. J. Johnson and W. C. Jankowski, in Carbon-I3 NMR Spectra (John Wiley, New York, 1972, p. 195; E. Breitmaier and W. Voelter, 13C NMR Spectroscopy (Verlag Chemie, Weinheim, 2nd edn, 1978). Paper 6/1998 ; Received 10th October, 1986

ISSN:0300-9599    

DOI:10.1039/F19888400041

出版商:RSC    

年代:1988

数据来源: RSC

摘要:

J . Chew. SOC., Faraday Trans. 1, 1988, 84(1), 57-64 Thermal Decomposition of Silver Squarate Andrew K. Galwey" and M. Abdel Aziz Mohamedt Chemistry Department, Queen's University, Berfast BT9 5AG, Northern Ireland Michael E. Brown Chemistry Department, Rhodes University, Grahamstown 6140, South Africa The kinetics of the thermal decomposition (473-510 K) of crystalline silver(1) squarate, Ag,C,O,, under reduced pressure in an accumulatory gas apparatus, have been studied for comparison with results obtained for the decompositions of the nickel(@ and copper(I1) salts. No melting was observed and the overall products of decomposition were solid silver particles in a carbonaceous residue, pseudomorphic with the original reactant crystallites, and gaseous CO and CO,. The isothermal a us.time curves were mainly deceleratory and approximated, for a < 0.5, to either the contracting-area or the contracting-volume rate equations, with an apparent activation energy of 190 & 8 kJ mol-l. Silver oxide powder was found to catalyse the decomposition and this, together with the presence of CO, in the gaseous products, led us to suggest that Ag,O is a reaction intermediate, which reacts further with product CO gas. This sequence of reactions in a solid-state decomposition and our failure to detect any recognisable reactant/product interface in electron microscopic studies of partially decomposed material are the central features of this study. The relationship between these results and those for silver oxalate and for nickel and copper squarates is discussed.The metal salts of organic acids provide a group of related substances which permit comparison of the effects of variations of either cation or of anion on thermal stability. The thermal decomposition of many of the simple carboxylates, such as formates and oxalates, have been investigated,' but the patterns of behaviour and proposed reaction mechanisms have not, as yet, been generally agreed. The thermal decomposition of silver squarate has not, to our knowledge, been previously investigated. This salt is particularly suitable for study because it can be prepared in the anhydrous form, so that anion decomposition is not preceded by a dehydration. In this it resembles silver oxalate2* and silver mal~nate,~ but contrasts with the reaction of nickel sq~arate,~ where the dehydration step immediately precedes decomposition and was identified as exerting control over the kinetics of anion breakdown.Silver squarate was also selected to permit comparisons to be made with the thermal reactions of copper squarate.'* ' The squarate anion contains no carboxyl group, and so its study is complementary to the more extensive work1 on the reactions of formates, oxalates and other carboxylates. Experimental Preparation of Silver Squarate Silver squarate was prepared by slowly mixing, at 320 K with continuous stirring, equal volumes of 0.045 mol dm-3 aqueous squaric acid and 0.090 mol dm-3 silver nitrate. The t Permanent address : Chemistry Department, Assiut University, Qena, Egypt. 5758 Thermal Decomposition of Silver Squarate white precipitate stood for 14 h at 298 K, was washed with distilled water and dried.Elemental analysis gave a composition approximating to the acid salt AgHC,04 H,O. The preparation was further treated by warming to 320 K with excess ethanol for 3 h, followed by filtration. The resultant yellow-green solid was dried and stored in the dark as a precaution against possible photolysis (e.g. silver oxalate2. 3). The analysis of the treated salt agreed with the theoretical composition for anhydrous silver squarate, Ag2C404* Kinetic Measurements Isothermal (k 1 K) kinetic studies were based on measurements at known times of the pressure of evolved gas in a constant volume, initially evacuated Pa) glass apparatus. A manually operated McLeod gauge and Baratron diaphragm gauge were used.Output from the Baratron gauge was recorded8 and values of pressures, times and temperatures were stored in a Sinclair Spectrum microcomputer for later kinetic analysis. The gaseous products (CO and CO + CO,) could be studied selectively by using cold traps (78 or 175 K) between the heated reactant and the gauge. Electron Microscopy Jeol 35CF and JSM840 scanning electron microscopes were used. Samples were precoated with a thin film of Au/Pd except for back-scattering studies in the JSM840 instrument. Samples of reactant, of product and of salt, partially decomposed to various known extents, were examined. Results Reaction Stoichiometry The reaction stoichiometry, determined from measurements of the masses of reactant, of the solid residue and of the pressures of gases evolved on completion of decomposition in the known-volume apparatus using either the 78 K trap (to measure CO only) or the 175 K trap (CO + CO,), was : Ag2C404(s) -+ 2Ag(s) + 1.44CO(g) + 0.96C02(g) + [C1.60,.61(s)~ The residual product was shown by X-ray diffraction to contain silver metal but no Ag20.Ag20 is known' to be reduced readily to Ag by CO well below the present reaction temperature. The carbon content of the residue, shown above in square brackets, as determined by combustion analysis, is in satisfactory agreement with that expected (by difference) from the reactant composition and the yields of the volatile products. Mass spectrometry of the gaseous products confirmed the evolution of CO and CO,. Smaller responses were observed at m/e ratios: 114 (possibly traces of squaric acid), 86 and 82 (unidentified), 80 (possibly C,02) and 58 (C20,H, or acid fragment).There was no indication of oxygen (0,) formation. This absence is strong evidence that the decomposition of silver(1) oxide, which yields molecular oxygen" and occurs' below the temperature of silver squarate decomposition (> 470 K), is not a contributory step in the present reaction. A check was made on whether the relative proportions of evolved CO and CO, varied systematically with time, temperature or reactant mass, by making duplicate, but otherwise identical kinetic measurements with either a 78 K or a 175 K trap. Results showed that reaction rates were not detectably different when CO, was allowed to accumulate as a gaseous product and when it was continuously and irreversibly condensed.After both types of experiment the final yields of both products wereJ . Chern. SOC., Faraday Trans. 1, Vol. 84, part 1 Plates 1 and 2 Plate 1. Scanning electron micrographs showing typical crystals of undecomposed reactant (a = 0.0) silver squarate (scale bar = 1.0 pm). Plate 2. Silver squarate crystallites decomposed in vacuum to a = 0.3 at 500 K (scale bar = 1.0 pm). A. K. Galwey, M. A. A. Mohamed and M. E. Brown (Facing p . 58)J. Chem. SOC., Faraday Trans. 1, Vol. 84, part 1 Plate 3 Plate 3. Completely decomposed (a = 1.00) crystallites of silver squarate (both scale bars = 1.0 pm). A. K. Galwey, M. A. A. Mohamed and M. E. BrownA . K. Galwey, M. A .A . Mohamed and M. E. Brown I 59 tlmin Fig. 1. Representative a us. time curves for the isothermal decomposition of silver squarate at different temperatures (K). identical. A further decomposition was completed at 493 K; the 175 K trap present was replaced by a 78 K trap, to condense product CO,, and the residual products were maintained at ca. 500 K for 15 h. The CO pressure was not reduced and the CO, pressure, (175 K trap), was also unchanged. The composition of the final gaseous products was thus not influenced by removal of CO, during or after reaction and the residual silver-carbon mixture does not catalyse the reaction 2co -b CO, + c. Electron Microscopy From electron microscopic studies at various extents of reaction, we could find no evidence of melting8* 11-15 or sintering during silver squarate decomposition. The final residual particles (fractional reaction, a = 1.00) were pseudomorphic with those of the original reactant.The appearance of typical reactant crystallites (a = 0.00) is shown in plate 1. Linear dimensions of individual particles were usually ca. 1 pm, although sometimes with one longer dimension. Surfaces were approximately planar, some included pits, and crystallite corners were usually rounded. Aggregates of many small crystallites were also present. Decomposition (a = 0.3 at 500 K) was accompanied by the appearance of small protuberances on the surfaces, often located at crystallite edges, (plate 2). The decomposition residue (a = 1.00) (plate 3) is comprised of rounded particles (identified from back-scattering electron microscopic measurements as silver crystallites) embedded in a coherent matrix [a carbonaceous polymeric (?) material], pseudomorphic with the original reactant crystallites.This residue is similar in appearance to that from silver malonate decomp~sition.~ Similar aggregation of metal as small rounded particles on an apparently immobile carbonaceous coherent matrix was also described8 for the decomposition of copper(1) malonate, where again the residual particles were pseudomorphic with the reactant crystallites.60 1 .o 0.8 ar 0.6 0.4 0.2 0 Thermal Decomposition of Silver Squarate 0 0 + ' 0 0 t 0 + 0 0 + o o + 0 t 0 0 + 0 + o o + o + o + o + o + o t o + o t: O + o 4* I 20 40 60 80 100 120 140 160 tlmin 0 Fig. 2. Effect of variation of sample mass on the thermal decomposition of silver squarate at 500K: 0, 1 1 ; +, 35mg.Kinetic Measurements Isothermal a us. time curves (473-510 K) for silver squarate decomposition (fig. 1) were predominantly deceleratory with no induction period. Initially (0 < a < 0.45) the reaction rate was almost constant (zero-order), but diminished when a > 0.45, reaching a minimum between 0.55 < a < 0.60. This was followed by an acceleratory process, which became deceleratory when a > 0.85. The intermediate acceleratory region was not observed at the lower end of the temperature range studied (490 K) and, even when present, kinetic behaviour was not very reproducible, depending to some extent on sample mass. Fig. 2 shows two representative a vs. time plots for identical decompositions at 500 K using different reactant masses (1 1 and 35 mg).The mid- reaction acceleratory process was absent when the lower reactant mass was used. This behaviour suggested that the kinetics of this reaction are pressure-dependent. Comparative experiments were thus made in which the total accumulated product gases were measured throughout, and in which the gaseous products were removed by evacuation (for 3 min) at a = 0.45 before the onset of the acceleratory reaction. This kinetic comparison is shown in fig. 3, in the form of a differential plot of Ap/At vs. time, where Ap is the pressure difference between consecutive readings made at constant time intervals, At, for two such experiments at 500 K. The later acceleratory process was eliminated, or very significantly reduced, following evacuation of product gases at a = 0.45.The kinetic data for the initial process (a < 0.5) obeyed, with equal acceptability," both the contracting-area and the contracting-volume equations.' The latter fitted the data over the wider range (fig. 4). The kinetic data could,also be satisfactorily expressed by the first-order expression for 0.30 < a < 0.70, but at higher values of a( > 0.7) obedience was poor. Reaction is thus best represented by a contracting-volume model.' It is concluded that decomposition is initiated at the original crystallite surfaces and that chemical changes proceed within a zone that progressively advances inwards. The activation energy for the first part of reaction (a < 0.5) was 190+8 kJ mol-' (483-508 K).A .K. Galwey, M. A . A . Mohamed and M. E. Brown aB, 00 c 0 0 0 0 0 0 0 0 OO OO 0 0 0 I I 40 80 1: tlmin 61 I Fig. 3. Plots of Ap/At vs. time for two experiments at 500 K: 0, product gases evacuated (3 min) at a = 0.45; 0, in the presence of the continued accumulation of the product gases. 0.2 ””, I I 3 v 3 0.1 0 20 40 60 80 100 120 140 160 tlmin Fig. 4. Test for conformity of the data shown in fig. 1 to the contracting-volume equation.] Pre-crushing the reactant appreciably increased the rate of decomposition. This is evidence that the reaction rate was influenced by the surface area of the original reactant crystallites. Mixing silver squarate with 10% (by mass) of silver metal powder did not change the kinetics of decomposition, the gaseous product yields (at a = 1.00), or the CO/CO, ratio.Addition, with crushing, of 10% (by mass) of silver oxide powder62 Thermal Decomposition of Silver Squarate (Ag,O) to the reactant accelerated the breakdown of silver squarate so strongly that reaction was too vigorous to permit kinetic measurements within the present temperature interval. Ag,O is thus identified as an excellent catalyst for the decomposition of silver squarate, and may also occur as an active reaction intermediate. Discussion Reaction Mechanism Discussion will be concerned with the two complementary aspects of behaviour that require''' l2 consideration in the formulation of the mechanism of a solid-state reaction : the reaction geometry and the chemistry of the changes occurring within the reaction zone.The absence of autocatalysis or of any acceleration of decomposition on mixing with silver, together with the absence from the electron micrographs of any recognizable interface, lead us to conclude that there is no well defined reactant-product-active contact zone within which the chemical changes occur preferentially. Our observations (plate 3) show that metallic particles tend to be generated and to grow on the external surfaces of the crystallites,' whereas the carbonaceous residue is effectively immobile, preserving the sizes and shapes of the original reactant particles. Such development of silver particles at sites remote from reaction is evidence that there is no catalytically active silver-silver squarate contact interface and that silver metal does not participate in the anion-breakdown step.This contrasts with the mechanism described for silver malonate., The kinetic obedience to the contracting-volume expression, taken with the increase in rate caused by reactant crushing, suggests that the decomposition zone progressively proceeds inwards from the original crystallite surfaces. No well demarcated interface, within which an autocatalytic chemical change was completed, could be recognized. The evidence was that silver metal migrated beyond the zone at which it was generated. Intracrystalline Chemistry In formulating a detailed reaction mechanism, we first eliminate several possible intermediates. (i) If the initial steps in reaction were electron transfers, this could yield cyclobutanetetraone : Ag,C,04 + 2Ag + c40,.This mechanism was excluded because there is no obvious or plausible reorganization whereby the (CO), intermediate could give the significant yields of product CO, observed. (ii) The unsaturated anion might plausibly rearrange to silver acetylide. This reaction was excluded because silver acetylide is unstable'' at reaction temperature and would decompose rapidly. (iii) It is most improbable that silver would be oxidized to Ag2+: Ag,C,O, + Ag i- Ag2+C40:- under the present predominantly reducing reaction conditions. The formation of CO, provides an important insight into the reaction mechanism. This requires transfer of oxygen between carbon atoms and can be most satisfactorily explained here through the intervention of Ag20.In contrast, CO, is formed in only a relatively small yield during nickel squarate decomp~sition,~ where the nickel residue is a more active heterogeneous catalyst than our silver product.A . K. Galwey, M. A . A . Mohamed and M. E. Brown 63 Our proposed reaction mechanism is as follows: [C,O,] + 2.5CO + [C,.,O,.,] (polymeric residue) [Ag,O] + CO + 2Ag + CO,. Cherall this is Ag,C,O, + 2Ag + CO, + 1.5CO + (c1.50,)5)n which is in satisfactory agreement with the stoichiometric data above and also explains the formation of ca. 1 mol of CO, per mol of salt decomposed. The reduction of Ag,O by CO is very rapid under reaction conditions, so that this intermediate is short- lived.' Reaction proceeding through two consecutive steps accounts for the absence of a reaction interface, since the silver metal product is not in direct contact with the solid reactant.Breakdown of the postulated intermediate, [C,O,], can be expected to yield CO, and the polymeric residue containing some oxygen. The mid-reaction acceleratory process, observed during decompositions with product- gas accumulation, is ascribed to changes in the participating chemical steps rather than interface geometry.'l? l2 We suggest that the product CO, within the residual carbonaceous layer, interacts directly with the undecomposed salt, or promotes its breakdown, when the prevailing pressure is sufficient. The overall chemical change is unaltered : Ag,C,O, + CO + 2Ag + CO, + [C,O,] but the first step and oxide reduction are accelerated. The subsequent decomposition of [C,O,] is unchanged because there is no CO involvement.It is improbable that this mid- reaction acceleratory process arises through self-heating. Moreover, while this pattern of changes of reaction rate is formally similar to Smith-Topley behaviour,' it seems improbable that the present products (CO and CO,, but not H,O) promote the textural changes that explain this characteristic behaviour of hydrates. Two central features of this proposed mechanism differ from the behaviour usually regarded as characteristic of solid-state decompositions or crystolysisl' reactions. These are the absence of an active reactant-product interface and the occurrence of a secondary reaction of an initial product (Ag,O) with a product gas (CO). Thus, as is often found in this field,'Y l9 the thermal reactions of silver squarate differ from those of related reactants containing common constituents, e.g. silver malonate, or nickel squarate., Although silver malonate decomposes in a similar temperature interval,, the generation and development of a reaction interface during this nucleation and growth process involves a quite different sequence of chemical changes and controls from those discussed here.This difference may be the consequence of the hydrogen in the reactant anion, CH,(CO,),, which permits an entirely different surface chemistry involving reactive chemisorbed intermediates. The decomposition of silver squarate also differs from that of silver oxalate,'-, which occurs in a lower temperature range and is a nucleation and growth process exhibiting a sigmoid a us.time curve. The first step in Ag,C,O, breakdown has been identified as electron transfer. This decomposition, however, yields Ag and CO, products only, thus reducing the possibility of surface deactivation of the silver metal by deposited carbon or chemisorbed CO, and consequently the residual product silver readily promotes anion breakdown. The reaction of silver squarate shows several points of dissimilarity with the decomposition of nickel squarate,, which, however, proceeds in a similar temperature range. The kinetics of decomposition of NiC,O, .2H,O are dominated by the precursor dehydration step and this salt yields a higher proportion of product CO. The reaction of silver squarate did, however, exhibit several points of similarity with copper 3 FAR I64 Thermal Decomposition of Silver Squarate squarate,'.? which will be considered in detail in the context of formulating a reaction mechanism for this latter compound, discussed in a forthcoming paper.? The authors thank Mr J.McCrae and his staff and Mr R. H. M. Cross for helpful advice in obtaining the electron micrographs. M.A.M. thanks the Egyptian Government and the ORS Award Scheme for Scholarships held during the period of this work. M.E.B. acknowledges financial support from the South African CSIR. References 1 M. E. Brown, D. Dollimore and A. K. Galwey, Comprehensive Chemical Kinetics, Vol. 22. Reactions in 2 A. Finch, P. W. M. Jacobs and F. C. Tompkins, J. Chem. SOC., 1954, 2053. 3 A. G. Leiga, J. Phys. Chem., 1966,70, 3254; 3260. 4 A. K. Galwey and M. A. Mohamed, J. Chem. SOC., Faraday Trans. I , 1985, 81, 2503. 5 A. K. Galwey and M. E. Brown, J. Chem. SOC., Faraday Trans. I , 1982, 78, 41 1. 6 M. E. Brown, A. K. Galwey and M. W. Beck, Zsr. J. Chem., 1982, 22,215. 7 A. K. Galwey, M. A. Mohamed, S. Rajam and M. E. Brown, to be published. 8 N. J. Carr and A. K. Galwey, Proc. R. SOC. London, Ser. A, 1986, 404, 101. 9 I. Nakamori, H. Nakamura, T. Hayano and S. Kagawa, Bull. Chem. SOC. Jpn, 1974,47, 1827. the Solid State (Elsevier, Amsterdam, 1980). 10 G. V. Malinin and Yu. M. Tolmachev, Russ. Chem. Rev., 1975,44, 392. 11 A. K. Galwey, Proc. 7th Znt. Con$ Thermal Analysis, Kingston, Ontario (Wiley, New York, 1982), 12 A. K. Galwey, Thermochim. Acta, 1985, 96, 259. 13 A. K. Galwey, R. Spinicci and G. G. T. Guarini, Proc. R. SOC. London, Ser. A, 1981, 378, 477. 14 A. K. Galwey and L. Poppl, Philos. Trans. R. SOC. London, Ser. A, 1984, 311, 159. 15 A. K. Galwey, L. Poppl and S. Rajam, J. Chem. SOC., Faraday Trans. I , 1983, 79, 2143. 16 M. E. Brown and A. K. Galwey, Thermochim. Acta, 1979, 29, 129. 17 J. D. McCowan, Trans. Faraday SOC., 1963, 59, 1860. 18 N. J. Carr and A. K. Galwey, Thermochim. Acta, 1984, 79, 323. 19 D. A. Young, Decomposition of Sofia3 (Pergamon, Oxford, 1966). p. 38. Paper 612314; Received 1st December, 1986

ISSN:0300-9599    

DOI:10.1039/F19888400057

出版商:RSC    

年代:1988

数据来源: RSC

摘要:

J . Chem. SOC., Furuduy Trans. 1, 1988, 84(1), 65-74 An Extended X-Ray Absorption Fine Structure Study of Heat-treated Cobalt Porphyrin Catalysts supported on Active Carbon Bob van Wingerden," J. A. Rob van Veen and Cees T. J. Mensch KoninklijkelShell-Laboratorium, Amsterdam (Shell Research B. V . ) , Badhuisweg 3, 1031 CM Amsterdam, The Netherlands Transition-metal chelates are good catalysts for the electrochemical reduction of oxygen, and their activity can be further improved by heat treatment. The present EXAFS study of 5,10,15,20-tetra-(p-chloro- phenyl)porphyrinatocobalt(III) supported on Norit BRX active carbon and heat-treated in dinitrogen at temperatures up to 850 "C shows (i) that upon adsorption the Co chelate remains intact, (ii) that after heating at 550 "C, when the oxygen-reduction activity is at its maximum, all Co is still present in its original (square-planar N4) environment and (iii) that at higher temperatures the CON, part decomposes to form, ultimately, metallic Co, with a concomitant decline in catalytic activity.Metallic Co dissolves in 4 mol dm-3 H,S04, but reacts in 4 mol dm-3 NaOH to give an oxyhydroxide, which is thought to have a capacity for reducing oxygen electrochemically. Carbon-supported transition-metal porphyrins and phthalocyanins are good catalysts for the electrochemical reduction of oxygen and the electrochemical oxidation of carbon monoxide in aqueous media.' Heat treatment in an inert atmosphere at temperatures from 500 to 800 "C generally improves the performance of these materials substantially, particularly in acidic sol~tions.~-~ Several studies have been devoted to identifying the nature of the active site in the heat-treated catalysts. According to Kaisheva et a1.,* a special highly active type of carbon is generated during heat treatment.Our previous result^,^ on the other hand, strongly argue in favour of the hypothesis that the central MeN, part of the chelate remains the site of electrocatalytic activity. This picture has now been challenged by Yeager et a1.l' on the basis of in-situ Mossbauer experiments (using 57CoTPP, TPP = tetraphenylporphyrin). They suggest that either some Co oxide, formed during the heat treatment, or a 'pyrolysed C-N surface', produced through reaction between the chelate and the subjacent carbon, a reaction thought to be catalysed by the transition metal, is responsible for the electrocatalytic activity of heat- treated catalysts.The fact that we opt for the CON, part remaining the active site is ascribed by Yeager et a1.l' to our having used samples which do not contain' the chelate in monolayer form. Their discussion gives rise to the following comments. (i) It is not likely that the difference between their conclusions and ours can be traced to our samples containing inultilayers or even crystallytes of the metal chelate : several of these materials have been prepared via equilibrium adsorption from solution, neat adsorption isotherms always being observed,2'4 while the CO, area of the carbon (Norit BRX), 650 m2 g-', is sufficiently large to accommodate all chelate molecules even if they are lying flat on the surface (which we think is the case6).(ii) The temperature they applied to the carbon- supported cobalt porphyrin (850 "C), is ca. 200 "C higher than that for which we observed optimum oxygen-reduction activity.6 (iii) The proposed formation of a cobalt oxide at 850 "C is curious, in that in previous studies the formation of metallic C O , ~ ~ or 65 3-266 EXAFS Study of a Cobalt Porphyrin Catalyst c =c / \ CI I I /rc .c $4 'c=c / I C l Fig. 1. Structure of 5,10,15,20-tetra(p-chlorophenyl)porphyrinatocobalt(III). metallic Fe in the case of carbon-supported iron porphyrins,' has been observed, as would be expected since the gaseous pyrolysis products (H,, CO and hydrocarbons) constitute a reducing atmosphere.(iv) We cannot subscribe to any model which does not associate the catalytic activity with the metal ions (in whatever form they are present), since we have found that, at least for heat-treated carbon-supported cobalt and iron porphyrins, the electrocatalytic activity is proportional to the metal content," and that the heat-treated Ir analogue is highly active for the electrochemical oxidation of CO,' a feature that is difficult to associate with a 'pyrolysed C-N surface'.'O However, we realize that the best foundation for the idea that the MeN, part persists in the heat-treated catalysts could be provided by EXAFS. We have applied this technique previously,'2 but although the data obtained were consistent with the above idea, they were not good enough to prove the point.It is for this reason that we have initiated a further EXAFS study, the results of which are presented in this paper. In contrast to our earlier work, we have employed, instead of CoTPP, 5,10,15,20-tetra(p- chlorophenyl)porphyrinatocobalt(III) [CoT(p-Cl)PP], which is thermally less stable. Experiment a1 Samples and Preparation The chelate, CoT(p-C1)PP (fig. l), was prepared as previously described' [sample (a)]. It was fixed on the Norit BRX support by slowly adding water to a slurry of the active carbon in a dimethylformamide solution of the chelate [sample (b)]. The loading amounted to 0.20 mg chelate per mg Norit BRX. Heat treatments were carried out in an 0,-free N, atmosphere; samples were kept for 1 h at the required temperature [sample (c) at 550 "C, ( d ) at 670 "C and (f) at 850 "C].A sample heat-treated at 670 "C was placed in contact with 4 mol dm-3 H,SO, to remove any metallic cobalt formed [sample ( e ) ] ; a sample heat-treated at 850 "C was placed in contact with 4 mol dm-3 NaOH, to allow a comparison with Yeager's Mossbauer results1' [sample (g)]. Oxygen-reduction activities were determined on Teflon- bonded floating electrodes ; details of the measuring procedure have been reported previ~usly.~B. van Wingerden, J . A . R. van Veen and C. T. J . Mensch 67 X-Ray Techniques X-Ray diffraction (X.r.d.) spectra were recorded on a Philips PW1050 vertical diffractometer equipped with an automatic divergence slit and graphite monochromator, using Cu K,,,, radiation. X-Ray absorption data were collected at room temperature around the Co K absorption edge at beam line 7.1 of the Daresbury synchrotron.The operating conditions during two runs comprised primary electron beam energies of 1.8 and 2 GeV and an energy resolution of the channel-cut Si(ll1) crystal of ca. 2 eV. For each sample at least two spectra were collected. Noise levels in the data for each scan were ca. 0.001. Samples were measured in air in the form of pressed self-supporting wafers of a suitable absorption thickness, Data reduction was performed using in-house software standard Fourier filter techniques and least-squares fitting in k-space13 with theoretical phase shifts and backscattering amplitudes from Teo and Lee.14 The principles of EXAFS may be found in a number of reviews and will not be discussed here,l3 These reviews also contain details on data treatment.In general we first performed single-shell fits on the separate peaks. The results were then refined in a multiple-shell fit for the first three nearest-neighbour atom shells in order to minimize any mutual interference effects on the data analysis. In particular, this can be of importance for the second- and third-shell peaks (see below), which show up as an incompletely resolved doublet peak in the radial distribution function (r.d.f.). From our experience with reference compounds and results from the literature we estimate that distances to the first-neighbour atom shell can be determined with an accuracy of 0.002 nm. The coordination number for the first and seciond shells can be derived from the amplitude of the corresponding EXAFS oscillation to within 20 %.Results Stationary polarization curves under oxygen in 4 mol dmA3 H,SO, were similar to those reported previously.6 Tafel slopes ranged from 40 mV decade-l for the untreated CoT(p- Cl)PP/Norit BRX sample to 55 mV decade-l for samples heat-treated at temperatures of 550 "C and higher. In fig. 2 we have plotted the specific oxygen-reduction activities as a function of heat-treatment temperature. Maximum activity is observed after a heat treatment at 550 "C. EXAFS Spectra The results of the fitting procedure on the EXAFS spectra are given in table 1. As an illustration of our discussion, the oscillatory EXAFS pattern after background removal and the corresponding raw radial distribution functions are shown in fig.3(a)-(g). All results shown are normalized to equal heights of the absorption edge. For convenience, in fig. 4 some results of three-shell fits in k-space are shown. In the r.d.f. of the base material [fig. 3(a)] we can distinguish a sharp peak corresponding to the nitrogen nearest neighbours, for which a distance of 0.196 nm is found. This is in good agreement with EXAFS data on low-spin cobalt porphyrins reported in the 1iterat~re.l~ At greater distances a doublet peak is observed. Best-fit results for this doublet were obtained with two carbon-neighbour shells at 0.299 and 0.335 nm. These results are in agreement with literature data on tetraphenyl- porphinatocobalt, which is expected to have a similar Co environment to our compound except for the Cl attached to the porphyrin rings.l' As this Co-Cl distance exceeds 0.4 nm, the presence of Cl will not influence the EXAFS results presented here.In principle, all the distances observed in the cobalt porphyrin structure (fig. 1) can be ascribed to atoms within one molecule. Coordination numbers for the first three shells68 EXAFS Study of a Cobalt Porphyrin Catalyst A 300 500 700 900 heat-treatment ternperaturel'c Fig. 2. Specific oxygen-reduction activity us. heat-treatment temperature for oxygen reduction in 8 mol dm-3 H,SO, at room temperature. a, CoT(p-CI)PP/Norit BRX; A, Norit BRX; 0, H,T(p-Cl)PP/Norit BRX. are expected to be 4 : 8 : 4. As can be seen from table 1, we find good agreement for the ratio of the amplitudes of the first and second shells.The third-shell amplitude is relatively high. However, too little is known about the uncertainty in the determination of amplitudes beyond the second shell to regard this as significant. To gain some impression of the accuracy in the determination of the first-shell coordination number we also measured the EXAFS spectrum of cobalt phthalocyanine. Here, too, Co is known to be surrounded by four nitrogen atoms. The resulting amplitude corresponding to the nitrogen nearest-neighbour shell is equal to that given in table 1 to within the experimental error claimed. When the porphyrin is deposited onto the carbon support [sample (b), fig. 3(b)], the EXAFS data clearly indicate that the structure of the unsupported porphyrin is retained. The distances found remain the same as in the unsupported compound.Although the amplitudes for the first- and second-neighbour shells are slightly higher, this is within the experimental accuracy and therefore cannot be regarded as a significant indication of any cobalt-support interaction, The ratio of the first- to the second-shell amplitudes is again ca. 1 : 2. For the Debye-Waller factor, which is related to the thermal vibration and/or structural disorder, a decrease is observed in the first-neighbour shell. This may be related to the fact that cobalt porphyrins are now isolated on the carbon surface instead of being incorporated in a larger crystalline structure. The lack of observation of a fixed cobalt-support interaction is in line with previously reported results on a similar cobalt porphyrin system and may probably be interpreted as an indication of random adsorption of the porphyrin molecules on the carbon surface.12 After heat treatment at 550 "C [sample (c)], the various distances between neighbours deduced from the analysis are still identical to those of the unsupported porphyrin. The amplitude of the first nitrogen neighbour shell also remains the same within experimental error, but for the second (and the third) shells a significant reduction in amplitude and hence coordination number is observed.About two out of the eight next-nearest carbonTable 1. Fourier filtering results for Co EXAFS data" ~~ first-neighbour shell second-neighbour shell sample A1 Rl XW,, z A2 R2 - G w . 2 Z . -~ ~ ____ ( a ) = CoT(p-Cl)PP (b) = ( a ) on Norit BRX (c) = (b) 550 "C (d) = (b) 650 "C (e) = (d)+H2S0, (f) = (b) 850 "C (g) = ( f ) + NaOH ( h ) = Co metal 1.38 1 S O 1.32 1.10 0.9 1 0.83 I .27 5.60 0.196 0.196 0.195 0.189 0.188 0.190 0.189 0.246 0.14 N 2.70 0.299 0.52 C 0.06 N 3.04 0.299 0.50 C 0.26 N 2.16 0.299 0.65 C 0.27 N 0.85 0.249 1.62 Co 0.48 N 0.93 0.295 1.58 C 0.90 N 2.81 0.247 1.28 Co 0.23 0 1.26 0.282 0.61 Co 1.07 Co - - - - Fourier ranges as follows : samples (a)-(g), 35-120 nm-' ; backtransform 0.1-0.33 nm ; sample (h), 35-1 20 nm-' ; backtransform 0.1 W.26 nm ; all with nm2) (twice the root mean- k3-weighting.A = amplitude of EXAFS oscillation; R = interatomic distance (nm); X,, = Debye-Waller factor square deviation of the distances); 2 = atom type used; % fit = least-squares agreement.third-neighbour shell Yo fit 1.93 0.335 0.77 C 12.4 2.31 0.335 1.16 C 12.0 1.43 0.333 0.79 C 12.6 9.6 0.78 0.336 0.65 C 7.6 7.7 4.5 2.9 ~~ A , R, Jkw., z - - - - - __ - - - - - - - - - - c b ? C 3 Q h30 150 0.8 0 0.0 30 150 0 0.8 Fig. 3. EXAFS oscillatory pattern k 3 ~ ( k ) , with k reciprocal wavevector (nm-') and corresponding raw radial distribution function. (a) CoT(p-Cl)PP, (b) (a) on Norit BRX, (c) (b) at 550 "C, (d) (b) at 670 "C, (d') 85 % (c) plus 15 % Co metal, (e) (d) plus H,SO,, cf) (b) at 850 "C, (g) df) plus NaOH, (h) Co metal. klnm-' Rlnm Rlnm klnm-' RlnmB. van Wingerden, J. A . R. van Veen and C. T. J. Mensch 10 0 - n y -10 71 ( a ) -10 5 30 1 klnm-' 10 Fig. 4. Three-shell fit pattern of k 3 ~ ( k ) for samples (a) and (b); (-) experimental, (---) fit. atoms seem to have been removed, together with a similar percentage of third-neighbour shell atoms.Simultaneously, we observe an increase in the first-shell Debye-Waller factor, which may be an indication of an increased disorder. Upon heating to 670 "C [sample (d)] the EXAFS pattern changes drastically. The distance of the nearest-neighbour shell changes to 0.189 nm, while the amplitude decreases by 20% compared with the unsupported porphyrin. The r.d.f. at larger distances becomes dominated by one peak at 0.249 nm, whereas the envelope of the corresponding EXAFS oscillation points to the presence of atoms heavier than C or N. Comparison with the reference spectra of a Co metal foil [fig. 3 (h) and table 11 shows that this second peak is representative of the formation of Co metal.Indeed, a line due to metallic Co appears in the X.r.d. spectrum (fig. 5). From the amplitude of the EXAFS oscillation related to this second peak one can estimate that ca. 15 O/O of the Co is present in the metallic state. This is in good agreement with the reduction in the amplitude of the Co-N peak and indicates that the remaining non-metallic Co atoms are still four- coordinated by nitrogen atoms. The presence of the dominant Co atom shell precludes a direct analysis of the amount of second- and third-shell carbon atoms. Nevertheless, in order to obtain information on the latter we simulated an r.d.f. resulting from 85% of Co atoms present as in sample (c) and 15% as metallic Co, by addition of the corresponding EXAFS oscillatory signals in k-space and subsequent Fourier trans- formation.The resulting r.d.f. is also given in fig. 3 ( d ) and shows that the number of carbon next-nearest neighbours must have decreased substantially in comparison with sample (c). In our opinion the decrease in Co-N is due to this further loss of part of the surrounding porphyrin rings, which allows a different binding of Co to N. After leaching with H,SO, the metallic Co has disappeared [sample (e), fig. 3(e)].72 EXAFS Study of a Cobalt Porphyrin Catalyst 0.60 0.40 A1 O . * O t P +a 0.04 OS6O 1 20.0 40.0 60.0 80.0 2810 0.80 '.OO 1 0.60 0.40 0.20 -5 0.00 >. E g 0.10 .- 0.80 0.00 I ' I I l l I l l 20.0 40.0 60.0 80.0 2810 Fig. 5. X.r.d. patterns of (a) Norit BRX support, (b) sample ( d ) heat-treated at 670 "C, (c) Co metal and ( d ) sample cf) heat-treated at 850 "C.(For X.r.d. a mixture of a and /3 cobalt powder, both with N , = 12 and R, = 0.251 nm, on a glass disc was used. The broad peak at 28 = 19" originates from the glass. For the other samples A1 sample holders were used. This gives rise to the A1 lines in their spectra.) Elemental analysis shows a decrease of ca. 25 YO in Co content, indicating that some Co is also probably leached out of the supported Co porphyrin. As the EXAFS results give the average surroundings of all the Co atoms present in the sample, one would expect the first-shell amplitude to be similar to that of samples (a)-(c) if the remaining Co atoms were all still four-coordinated by nitrogen. This is not the case, however: the amplitude of the first shell has further decreased by ca.35% compared with the unsupported porphyrin. This may indicate a loss of nitrogen nearest-neighbour atoms, but it is more likely that upon the leaching treatment some Co from the destroyed porphyrin still remains present on the support in an as yet unexplained but random way. These Co atoms would not contribute to the EXAFS signal but would still affect the normalization, hence causing an overall reduction in the amplitudes. This may also be reflected in a further increase in the first-shell Debye-Waller factor. In any case, the ratio between the first- and second-shell amplitudes is again much lower than in sample (c), in line with our findings for the unleached sample. In fig.3 0 we show the r.d.f. for sample df) treated at 850 "C. Here too we observe a short Co-N bond length and the presence of Co metal, the latter also being evident from the X.r.d. pattern (fig. 5). If we normalize on the amplitude of the Co metal foil reference sample, it is estimated that ca. 50% of the Co present is in the metallic state. Within the experimental accuracy this is in good agreement with the reduction in Co-N amplitude compared with the unsupported porphyrin. Furthermore, it is noticeable that in comparison with the 670 "C treated sample there is a further increase in the Debye-Waller factor. Again no indications can be obtained of the presence of any carbon neighbours beyond the first nitrogen coordination shell. Finally, in the sample treated at 850 "C, and afterwards with NaOH [fig.3(g)], weB. van Wingerden, J . A . R. van Veen and C. T. J . Mensch 73 observe one prominent neighbour peak at a short distance and one at a longer distance. The latter can be nicely fitted by Co at 0.282 nm. This indicates that some cobalt(hydr)oxide has been formed. CoO(0H) seems a likely candidate, in view of the similarity of its Co-0 and Co-Co neighbour distances to our re~u1ts.l~ However, a CON, species may also still be present, which complicates matters. A full analysis has not been attempted. Discussion From the results described the following pyrolysis reaction sequence can be deduced : 550 'C 670°C 67OOC [COP], + [COP], - [. * .I - co, 850 "C where COP denotes the carbon-supported cobalt porphyrin and the subscripts p, c and m stand for physisorbed, chemisorbed and metallic, respectively.However, note that at no stage is pyrolysis complete: [COP], is found to persist even after being heated at 850 "C. The chelate is adsorbed intact. A heat treatment at 550 "C does not disturb the central CON, unit, but induces some changes in the remoter parts of the molecule. This is consistent with our view, derived from various other experimental observations [cf. Introduction in ref. (9)] that in the first instance only (some of) the meso-carbon atoms are removed from the chelate, which in turn leads to the reaction of the affected pyrrole fragments with the subjacent carbon: the chelate is now chemisorbed. It is evident, therefore, that the oxygen-reduction activity is at a maximum (cf.fig. 2 ) while all of the Co and N is still present as the original CON, part of the sorbed chelate. Upon heating at higher temperatures, a progressive decomposition of the chelate is observed with concomitant formation of metallic Co, the latter species being the expected product (cf. Introduction). From the EXAFS data on the sample heat-treated at 670 "C, it seems that there is an intermediate stage between [COP], and Corn, indicated by the dots in reaction sequence (i) (dots are used because we have no idea as to the appearance of this stage). It appears as CON, in the EXAFS spectrum of the 670 "C sample as such, but it is only present in rather irregular surroundings in the sample treated with 4 mol dm-3 H,SO,. The progressive breakdown of the chelate molecules as the temperature increases is clearly shown by the continuous decrease in the amplitude of the peaks due to second- and third-shell neighbours.Note that with increasing decomposition of the cobalt porphyrin, the oxygen-reduction activity decreases. Since the metallic cobalt dissolves in 4 mol dmP3 H,SO,, it cannot contribute to the oxygen-reduction activity of the heat-treated catalyst in that electrolyte. In 4 mol dmP3 NaOH the situation is different : some cobalt(hydr)oxide, e.g. CoO(OH), is formed, and it would appear that it is the Mossbauer spectrum of this species that has been observed by Yeager et a1.l' In view of the known oxygen-reduction activity of such Co-containing materials as COA~,O,/C,~* NiCo,O," and La,-,Sr,Co0320 in alkaline media, it is quite probable that this cobalt(hydr)oxide is active as well.Yeager et a1.l' also observ d a substantial activity for 0, reduction for CoTMPP/Vulcan XC-72 heat-treated at 850 "C, in 85 % H3P0, at 100 "C. In our opinion this activity is due to the CON, species still left in the pyrolysed material: our EXAFS data indicate the continued presence of this species, even after pyrolysis at 850 "C, and fig. 1 shows that although the oxygen- reduction activity (in 4 mol dm-3 H,SO, at room temperature) is less than optimal, it is still very much higher than the carbon-only activity. Finally, note that the X-ray diffractograms of the heat-treated CoT(p-C1)PPINorit RRX samples (fig. 5 ) do not show the presence of the special type of carbon which Gamburzev et a1.' consider to be the actual active phase in the pyrolysed material.The a74 EXAFS Study of a Cobalt Porphyrin Catalyst 0 = 26" peak is entirely due to B-Co, the intensity paralleling the evolution of metallic Co observed in EXAFS. Conclusions (i) The central CON, part of the porphyrin system in CoT(p-C1)PPINorit-BRX materials remains intact upon heat treatments at temperatures up to 550 "C. At higher temperatures the chelate starts to break down, with concomitant formation of metallic Co. (ii) Maximum oxygen-reduction activity in 4 mol dm-3 H,SO, is observed for the sample pyrolysed at 550 "C, i.e. for a material in which the central part of all the adsorbed chelate molecules is still intact. Pyrolysis at higher temperatures leads to less active materials, in parallel to the decomposition of the chelate. Therefore, the conclusion that the CON, part is responsible for the observed oxygen-reduction activity in 4 mol dm-3 H,SO, seems inescapable.(iii) Metallic cobalt dissolves in 4 mol dm-3 H,SO,, but in contact with (air-containing) 4 mol dm-3 NaOH it is transformed into a cobalt(hydr)oxide species, which is expected to have some activity for 0, reduction in an alkaline medium. References 1 J. A. R. van Veen and J. F. van Baar, Rev. Inorg. Chem., 1982, 4, 293. 2 V. S. Bagotzky, M. R. Tarasevich, K. A. Radyushkina, 0. A. Levina and S. I. Andrusyova, J. Power 3 K. Wiesener and A. Furhmann, Z. Phys. Chem. Leipzig, 1980, 261, 411. 4 H. Jahnke, M. Schonborn and G. Zimmermann, Ado. Chem. Rex, 1976, 61, 133. 5 J. A. R. van Veen and C.Visser, Electrochim. Acta, 1979, 24, 921. 6 J. A. R. van Veen, J. F. van Baar, C. J. Kroese, J. G. F. Coolegeen, N. de Wit and H. A. Colijn, Ber. Bunsenges. Phys. Chem., 1981, 85, 693. 7 J. F. van Baar, J. A. R. van Veen and N. de Wit, Electrochim. Acta, 1982, 27, 57; J. F. van Baar, J. A. R. van Veen, J. M. van der Eijk, Th. J. Peters and N. de Wit, Electrochim. Acta, 1982, 27, 1315. 8 A. Kaisheva, S. Gamburtzev and I. Iliev, Sou. Electrochem., 1982, 18, 127; G. Gruenig, K. Wiesener, A. Kaisheva, S. Gamburtsev and I. Iliev, Sou. Electrochem., 1983, 19, 1408; G. Gruenig, K. Wiesener, S. Gamburzev, I. Iliev and A. Kaisheva, J. Electroanal. Chem. 1983, 159, 155; A. Fuhrmann, K. Wiesener, 1. Iliev, S. Gamburzev and A. Kaisheva, J. Power Sources, 1981, 6, 69. 9 J. A. R. van Veen, J. F. van Baar and C. J. Kroese, J. Chem. Soc., Faraday Trans. I , 1981, 77, 2827. 10 D. A. Scherson, S. L. Gupta, C. Fierro, E. B. Yeager, M. E. Kordesch, J. Eldridge, R. W. Hoffman 11 J. A. R. van Veen and H. A. Colijn, Ber. Bunsenges. Phys. Chem., 198 1, 85, 700. 12 R. W. Joyner, J. A. R. van Veen and W. M. H. Sachtler, J. Chem. Soc., Faraday Trans. 1, 1982, 78, 13 P. A. Lee, P. H. Citrin, P. Eisenberger and B. M. Kincaid, Rev. Mod. Phys., 1981, 53, 769. 14 Boon-Keng Teo and P. A. Lee, J. Am. Chem. Soc., 1979, 101, 2815. 15 A. Michalowicz, Now. J. Chim., 1982, 6, 79. 16 P. Madura and W. R. Scheidt, Inorg. Chem., 1976, 15, 3182. 17 S. J. Gurman, J. Muter. Sci., 1982, 17, 1541. 18 K. V. Kordesch, in Handbook of Fuel Cell Technology, ed. C. Berger (Prentice-Hall, Englewood Cliffs, 19 W. J. King and A. C. C. Tseung, Electrochim. Acta, 1974, 19,485; 493. 20 G. Bronoel, J. C. Grenier and J. Reby, Electrochim. Acta, 1980, 25, 1015. Sources, 1977178, 2, 233. and J. Blue, Electrochim. Acta, 1983, 9, 1205. 1021. NJ, 1968). Paper 612330; Received 2nd December, 1986

ISSN:0300-9599    

DOI:10.1039/F19888400065

出版商:RSC    

年代:1988

数据来源: RSC

摘要:

J. Chem. SOC., Furuduy Trans. I, 1988, 84(1), 75-85 Effect of Solvent on the Reactions of Coordination Complexes Part 2.-Kinetics of Solvolysis of cis-(Chloro)(imidazole)bis(ethylenediamine)- cobalt (111) and cis-( C hloro)( benzimidazole) bis( e t hy1enediamine)co bal t ( 111) in Methanol-Water and Ethylene Glycol-Water Media Anadi C. Dash* and Neelamadhab Dash Department of Chemistry, Utkal University, Bhubaneswar 751 004, India The kinetics of solvolysis of cis-(chloro)(imidazole)bis(ethylenediamine)- cobalt( 111) and cis-(chloro)(benzimidazole)bis(ethylenediamine)cobalt( 111) have been investigated in aqueous methanol (MeOH) and aqueous ethylene glycol (EG) media (040% by weight of MeOH or EG) at 45-64.7 "C. The logarithm of the pseudo-first-order rate constants for MeOH-water media exhibits linear dependence with the reciprocal of the bulk dielectric constant (oil), the mole fraction of MeOH (XMeOH) and the solvent ionizing power Y ( q.AdC1) as determined by the solvolysis rates of 1-adamantyl chloride.Similar plots (logk",,, us. x,, or oil) for EG-water media are non-linear. It is evident that the solvation phenomenon plays dominant role and the rate of solvolysis is mediated by the dual solvent vectors, the overall acidity and basicity of the solvent mixtures. The relative transfer free-energy calculations indicate that the mixed solvent media exert more destabilizing effect on the transition state as compared to the initial state. The activation enthalpy and entropy vs. Xorg (where Xorg is the mole fraction of the organic solvent component) plots display maxima and minima indicating that the solvent structural changes play significant role in the activation process.The activation free energy at a given temperature, however, increases only marginally and linearly with increasing Xorg. The mutual compensatory effect of activation enthalpy and entropy on the activation free energy is in keeping with the fact that the perturbations of the reaction zone and the solvent network remain approximately proportional to each other with increasing Xorg so that the isodelphic and the lyodelphic components of AH* and AS* correlate well with each other. En a previous paper1 the kinetics of solvolysis of the cis-[C~(en),(bzmH)Br]~+ (where bzmH = benzimidazole) was reported in methanol-water media over an extended range of solvent composition (0-80% methanol by weight) and at 35-55 "C.The rate data (logkz,,) exhibited a marked departure from a linear correlation with the reciprocal of the bulk dielectric constant of the medium and the Grunwald-Winstein parameter, Y, or the revised Y values (for 1-adamantyl chloride) reported by Bentley and Carter., Attempts to correlate logkz,, with the mole fraction of MeOH (XMeOH) were most successful. Plots of logk",,, us. X,,,, were excellent straight lines at 50 and 55 "C; however, increasiung positive deviation from linearity was noted at lower temperatures. 'These facts led us to believe that the specific solvation effects presumably become less significant with increasing temperature. The trend in the variation of the activation enthalpy and entropy with XMeOH indicated that the solvent cosphere of the substrates in the initial state (i.s.) and the transition state (t.s.) exerts a significant modulating effect on the rate and the thermodynamic parameters of the solvolysis process.For the aqueous alcohol system it is normally expected that the interaction of the 7576 Reactions of Coordination Complexes solvent with the substrate both in the initial state (is.) and the transition state (t.s.) will very much depend upon the acidity and basicity of the leaving groups as well as of the solvent. In this context it was felt worthwhile to extend the solvent-effect studies to other suitable cobalt(1n) substrates which can interact favourably with the solvent systems. In this paper we present some of our findings on the kinetics of solvolysis of the cis- (chloro)(imidazole)bis(ethylenediamine)cobalt( HI) and the cis-(chloro)(benzimidazole)- bis(ethylenediamine)cobalt(m) ions in methanol-water and ethylene glycol-water media. Experimental The cis-(chloro)(imidazole)bis(ethylenediamine)cobalt(rII) diperchlorate and the cis- (chloro)(benzimidazole)bis(ethylenediamine)cobalt( 111) diperchlorate were prepared and purified as described e a ~ l i e r .~ ' ~ The purities of the samples were checked by analysis of Co and Cl- which agreed to better than fO.l YO of the calculated values for the respectively. AnalaR methanol and ethylene glycol (EG) were further dried over 4 A molecular sieve and distilled; the middle fraction was collected.Gas chromatography using an Aimil Nucon model 5700 gas chromatograph did not reveal the presence of other impurities. Solvent mixtures were prepared by weight in the usual way. Spectro- photometric measurements were made using an LKB Biochrom Ultrospec I1 spectrophotometer or a u.v.-visible spectrophotometer manufactured by the Electronic Corporation of India Ltd. The procedure for following the kinetics of solvolysis of the complexes and determining the observed pseudo-first-order rate constants ( e b s ) has been described in detail in our earlier paper.l [Co(C2H,N2)2(C3N2H4)c11(c104)2 and [Co(C2H,N2)2(C,N2H,)CII(CIO,),, Results and Discussion The ligand-field bands of the chloro complexes are influenced little by the organic solvent components of the mixed solvent media.Successive spectral scans (400 < A/nm < 580) during the solvolysis of the ci~-[Co(en)~(bzmH)Cl]~+ in a 50 % MeOH-water mixture displayed isosbestic points at 505 and 427 nm with E = 77.5 and 27.5 dm3 mol-' cm-', respectively, which compare well both in position and intensity with the aquation reaction of the same carried out in a fully aqueous medium [A/nm ( E dm3 mol-' cm-') : 500 (79), 429 (28)]. However, at a relatively high percentage of EG and MeOH and at long reaction times the isosbestic points shifted, thereby indicating that both the solvent components (H,O/EG or MeOH) of the medium compete for the cobalt(rI1) centre with water predominating during the solvolysis reaction depicted in eqn (1) : cis- [Co( en),( B)Cl] + % cis- [ Co( en) 2( B)S] 3+ + Cl - (1) where S = solvent.The rate data at various solvent compositions are presented in tables 1-3. The activation parameters, calculated by the least-squares fitting of the rate data' to the transition-state equation : In kzbs = (In k / h + AS* / R) - (1 O-3AH' / R ) X (1 03/ T ) (2) are collected in table 4. Variation of Rate with Solvent Composition The observed pseudo-first-order rate constant (k:bs) decreases with increasing amount of organic solvent (see tables 1-3). Also, the plots of logk:,, against the reciprocal of theA . C. Dash and N . Dash 77 Table 1. Rate data for solvolysis of cis-[Co(en),BC1I2+ (B = imH, bzmH) in methanol-water mixtures: [complex], = (2-3) x mol dm-3; [HCIO,], = 0.01 mol dm-3 k&J l 0-5 s-' a, MeOH (WtYo) XMeOH 45.0f0.1 "C 50.2f0.1 "C 55.1 kO.1 "C 59.8f0.1 "C 0.00 5.00 10.0 20.0 30.0 40.0 50.0 60.0 70.0 80.0 0.00 0.0287 0.0588 0.1233 0.1942 0.2727 0.3600 0.4576 0.5675 0.6923 1.69 f 0.05 1.58 & 0.05 1.39 f 0.08 1.24 & 0.09 0.97 f 0.03 0.69 f 0.04 0.53 f 0.03 0.44 f 0.02 0.25 f 0.02 0.16 f 0.07 3.20 f 0.13 (1.26 f 0.11) 2.95 fO.10 (1.25 f 0 .11) 2.60 f 0.12 (1.02 f 0.02) 2.24 f 0.06 (0.83 f 0.01) 1.80 f 0.04 (0.63 f 0.04) 1.15 f 0.02 (0.50 f 0.03) 0.87 f0.02 (0.34 f 0.04) 0.72 f0.05 (0.30 f 0.03) 0.42 _+ 0.05 0.31 f0.05 5.53 k0.18 4.87 kO.11 4.32 f 0.23 3.54 k 0.04 2.63 f 0.07 1.72 f 0.05 1.24 f 0.04 1.01 f0.03 0.63 f 0.04 0.43 f 0.02 8.98 & 0.28 7.55 f 0.24 6.93 f 0.22 5.35 f 0.20 3.65 f 0 . 10 2.58 & 0.09 1.80 f 0.05 1.35 f 0.04 0.80 f 0.05 0.54 f 0.05 a Values in parentheses at 50.2 "C are for B = imH; all other values are for B = bzmH.Mean value and the error quoted as standard deviation were calculated from 10-1 5 individual values of ktbs from duplicate or triplicate runs, respectively, at each solvent composition. Table 2. Rate data for solvolysis of cis-[Co(en),(imH)Cl]"+ in ethyleneglycol-water mixtures : [complex], = (2.74.0) x mol dmP3; [HCIO,], = 0.01 mol dm-3 gbs/ 10-5 s-1 a EG (wt Yo) X,, 50.0f0.1 "C 54.8kO.l "C 59.7f0.1 "C 64.7fO.l "C 0.00 5.00 10.0 20.0 30.0 40.0 50.0 60.0 70.0 80.0 ~ ~ _ _ 0.00 0.015 0.03 14 0.0677 0.1 106 0.1622 0.2250 0.3034 0.4038 0.5373 1.2 1 f 0.03 1.14 f 0.02 1.07f0.01 0.90 4 0.02 0.78 f 0.03 0.62 f 0.02 0.54 f 0.02 0.44 f 0.03 0.36 f 0.02 0.27 f 0.01 2.00 f 0.05 1.86 f 0.04 1.83 f 0.06 1.60 _+ 0.04 1.38 fO.01 1.12 f0.04 0.88 f 0.08 0.83 & 0.04 0.60 f 0.04 0.38 k0.06 3.64 f 0.12 3.38 f 0.09 3.14 f 0.04 2.74 f 0.09 2.37 f0.06 1.97 f 0.03 1.72 f 0.05 1.38 f 0.03 1 .1 8 f 0.04 0.88 f 0.04 5.51 f 0.15 5.48 f 0.1 1 5.17f0.18 4.84f0.16 3.78f0.10 3.28 f 0.03 2.70f0.13 2.23 f 0.07 1.65 f 0.04 1.26 & 0.03 a See footnote (6) of table 1. bulk dielectric constant of the medium5 (Dil) are reasonably good straight lines (correlation coefficient = 0.99 1-0.996) for the MeOH-water medium ; however, such plots for EG-water tend to be curved [fig. 1 (a)]. The plots of logk& against the revised solvent ionising power of MeOH-H,O media reported by Bentley and Carter,2b q.AdCI (1-adamantyl chloride scale), are reasonably linear [fig.1 (b)] at 50 "C (correlation coefficient = 0.997) with slope = 0.23 and 0.25 for imidazole and benzimidazole complexes, respectively, which agree well with the values of the slopes of analogous plots for the aquation of the halogenoamine cobalt(n1) complexes in relatively water-rich alcohol-water media.', The plots of log k& against the mole fraction of MeOH (XMeoH) also generate straight lines for the benzimidazole complex (correlation coefficient = 0.997-0.998) with slopes = - 1.42 & 0.04 to - 1.78 0.03 at 45 to 59.8 "C. Similar plots for imidazole complex were also linear for the MeOH-water and EG-water media at 50 and 64.8 "C with slopes of - 1.46 0.03 and - 1.28 f 0.03, respectively (correlation78 Reactions of Coordination Complexes Table 3.Rate data for solvolysis of cis-[Co(en),(b~mH)Cl]~+ in ethylene glycol-water mixtures : [complex], = 2.0 x mol dm-s; [HCIO,], = 0.01 mol dm-3 gb,/ 10-5 s-1" EG (wt%) X,, 45.0f0.1 "C 50.0f0.1 "C 54.9fO.l "C 64.7f0.1 "C 0.00 5.00 10.0 20.0 30.0 40.0 50.0 60.0 70.0 80.0 90.0 0.00 0.015 0.03 14 0.0677 0.1 106 0.1622 0.2250 0.3034 0.4038 0.5373 0.7232 1.69 f 0.02 1.64 f 0.05 1.46 f 0.09 1.3 1 f 0.09 1.20 f 0.05 1.01 f0.09 0.92 f 0.08 0.84 f 0.05 0.62 f 0.06 0.45 f 0.07 3.62 f 0.12 2.88 f 0.05 2.62 f 0.09 2.38 f 0.09 2.09 f 0.1 1 1.62 f 0.07 1.45 & 0.03 1.15 f 0.07 0.95 f 0.05 0.72 f 0.03 5.57 f 0.19 4.89 f 0.07 4.53f0.14 3.97 f 0.23 3.17k0.11 2.70 f 0.06 1.96 f 0.10 1.59f0.11 1.32 f 0.06 0.95 f 0.05 15.20 f 0.5 14.00 f 0.7 13.10f 0.3 1 1 .SO f 0.6 9.18 f 0.35 8.04 f 0.47 6.09 f 0.19 4.61 f0.21 3.40 f 0.10 2.36 f 0.09 1.55 f 0.04 ~~ ~~~ a See footnote (b) of table 1.Table 4. Activation enthalpy, entropy and free-energy data for the solvolysis of cis-[Co(en),BC1I2+ (B = imH, bmH) in various MeOH-water and ethylene glycol-water mixtures MeOHb EGb orga AH* AS* AGkC AH* A S AG*" (wt %) /kJ mol-' /J K-l mol-' /kJ mol-' /kJ mol-' /J K-' mol-1 /kJ mol-' 0.00 5.00 10.0 20.0 30.0 40.0 50.0 60.0 70.0 80.0 96.6f 1.1 90.6 f 2.9 91.6f1.9 81.5 f 2.9 73.9 f 7.3 73.3 f2.0 70.7 f 2.9 63.1 f 3.9 65.6 f 6.4 52.0 f 9.9 -33f3 -52f9 -5Of6 -83f9 - 108 f 22 -113f6 - 124f9 - 149f 12 - 146f 19 - 190f30 107.1 f 0.1 (109.7 f 0.2) 107.5 f 0.1 (109.7 f 0.2) 107.8 f 0.1 108.2 f 0.1 (110.8f0.0) 108.8 f 0.1 (111.6f0.1) 110.0 f 0.1 (1 12.2 f 0.2) 110.8 f O .1 ( 1 1 3.2 f 0.2) 11 1.3 f0.2 (1 1 3.5 f 0.2) 112.7f0.3 (110.2f0.0) 113.5 f0.4 96.6f 1.1 (95.3 f4.2) 93.4f 5.6 (96.7 f 3.0) 97.7f 1.1 (97.4f2.3) 95.0 f 1.3 (106.0 f 2.4) 89.2 f4.1 (94.8 f 3.4) 93.5 f 5.0 (1OO.Of 1.7) 83.7 f 6.8 (99.4 f 4.8) 76.9 f 7.4 (94.0 f 4.5) 77.5 f 4.5 (93.3 f 9.5) 78.0 f 5.7 (92.8 f 6.7) -33f3 (-45 f 13) -44f 17 (-41 +9) -31f3 (-40f7) -4Of4 (- 14f9) -60f 12 (-50+_10) -49 f 15 (-36f5) -80f20 (-39f 15) - 102f22 (-57f 13) - 102+ 13 (-61 f28) -103f17 ( - 65 f 20) 107.1 f 0.1 (110.3f0.1) 107.5 f 0.0 107.7f0.1 108.0 f 0.1 (1 10.6 f 0.1) 108.3f0.1 (1 1 1 .o f 0.1) 109.0 f 0.1 (1 11.6 f0. 1) 109.3f0.1 1 09.9 f 0.1 (112.5f0.2) 110.4fO.l (113.1 fO.1) 111.2f0.1 (1 13.8 f0. 1) (1 10.0 f 0.0) (1 10.1 f 0.0) (1 12.0 f0. 1) a Organic solvent component.Parenthesized values are for the imidazole complex; all other values are for the benzimidazole complex. At 323.2 K. coefficient = 0.998). At all other temperatures the logk:,, vs. X,, plots for both the complexes are, however, non-linear. Typical plots are presented in fig. 2. Such non-linear plots are indicative of the specific solvation effects of EG-water medium on the solvolysis rates. However, we are led to believe that the substrate, cis-[Co(en),BCll2+ (where B = imH or bzmH), presumably does not discriminate between the solventA . C. Dash and N . Dash '1 -AdCl 4.0 3.0 2.0 1 .o 0.4 79 0 0.5 h " I v1 ---- 2 Y, w + - 1.0 1. 5 I 1 I 0 1.0 2.0 3.0 4 .O 1 O2 ID, Fig. 1. (a) Plots of 6+log(kEb,/s-') us. 102/D, at 50 "C for (1) the imidazole complex and (2) the benzimidazole complex; 0, MeOH-water; A,' EG-water.(b) Plots of 6 +log (ktbs/s-l) us. q-AdCI (MeOH-water) plots at 50 "C for (3) the imidazole complex and (4) the benzimidazole complex. components, MeOH or H,O, in its solvation envelope or skin phase.' Both the initial state and the transition state seem to be solvated by both solvent components without being selective for either MeOH or H20. It is most likely that the solvation of the substrates depends upon the dual solvent vectors, i.e. the acidity (A,) and basicity (B,) parametersa of the solvent mixtures, which for MeOH-water vary linearly with XIMeOH over the entire composition range studied [see eqn (4) and ( 5 ) of ref. (l)]. The same analysis is also applicable to EG-water medium, for which the assigned values of the acidity and basicity parameters of EG and H 2 0 (AEG = 0.78, AHIS) = 1.0; BEG = 0.84, BHZO = 1 .O)* also predict negative slopes of the straight h e plot of log kzbs us.XEG [see eqn (4H7) of ref. (1)) The structural perturbations in this solvent system, however, tend to modulate the solvent cosphere of the substrates, so that the linearity in the log kzbs us. X,, plot is not maintained over the entire composition range studied. The biphasic nature of the linear plots of log k& us. XFG at higher temperatures for the benzimidazole complex (see fig. 2) is then reconciled with the non-constancy of the parameters I, and I , of the solvent-effect relationship : [c is a constant ; see eqn (7) of ref. (l)] over the entire composition range. Interestingly, I , and Z2 are virtually constant for the imidazole complex in the XE, range = 0-0.537 at 64.7 "C, thus reflecting profound thermal effects on the structural perturbations in the EG-water mixtures and the disappearance of the specific effect of solvation on the solvolysis rate at relatively high temperatures.An excellent linear correlation between log k;,,(RCl2+) and log k",,,(RBr2+) {R = [Co(en),(bzmH)I3+, see ref. (1) for data for the bromo complex} over the whole80 Reactions of Coordination Complexes Xorg Fig. 2. Plots of 6+log(k",,,/s-') us. Xorg (where Xors is the mole fraction of the organic solvent component) for the solvolysis of ci~-[Co(en),BCl)]~+ [B = benzimidazole for 1 (64.7 "C), 4 and 6 (50 "C); B = imidazole for 2 (64.7 "C), 3 and 5 (50 "C)].I + z+ ("obs), t (enIzCo (BICI Is Scheme 1. Dissociative mode range of solvent compositions, 0 < XMeOH < 0.6923 and at 50 "C with slope = 0.93 kO.03 (correlation coefficient = 0.996) can be taken to indicate that there is no change in the mechanism of solvolysis with change in solvent composition. The familiar dissociative interchange me~hanism(I,),~ in which the Co-Cl bond is stretched to its limit in the transition state may be depicted as in scheme 1. Using the conventional thermodynamic cycle it can be shown that Kbs and ktbs (the rate constant in the fully aqueous medium) are related to the free energy of transfer of the transition state, [AGt(t.s.)](,,,,, and the initial state, [AG,(i.s.)],,+,,, as given by: logk;,, = logk,",, -(1/2.303~~)[AG,(t.s.)-AGt(i.s.)]~,,,, (3) where transfer of species is assumed to take place from water (w) to the mixed solvent (s).The non-electrostatic component of the relative transfer free-energy term, which is supposed to include all changes in the free energy that result in structural changes in the solvent when the initial state and the transition state of a given complex are transferred from water to the solvent mixtures, do not appear to be sensitive to the solvent composition of the MeOH-water mixture. It is the electrostatic component of the relative transfer free energy which varies linearly (having a negative slope) with the reciprocal of the bulk dielectric constant of the MeOH-water mixture [see fig. 1 (a)]. This is consistent with the fact that the ionization of the Co-Cl bond plays a dominant roleA .C. Dash and N . Dash 81 Table 5. Calculated values of the relative transfer thermodynamic functions at 323.2 K org" (wt%) MeOH EG MeOH EG MeOH EG 5.00 10.0 20.0 30.0 40.0 50.0 60.0 70.0 80.0 0.2 f 0.1 (0.0 f 0.3) 0.5fO.1 (0.6 f 0.2) 0.9 f 0.1 (1.1 f0.2) 1.5f0.1 (1.9 + 0.2) 2.7 f 0.1 (2.5 f 0.2) 3.5k0.1 (3.5 k0.2) 4.0 f 0.2 (3.9f0.2) 5.4 f 0.3 6.3 f 0.4 0.6 f 0.1 - 6 f 3 (0.2 f 0.1) 0.9 f 0.1 - 5 f 2 1.1 fO.1 -15f3 (0.3 f 0.1) (0.8 f 0.1) 1.5 f O . 1 -23f2 (1.2f0.1) 2.2 f 0.1 -23f3 (1.8 & 0.1) 2.5 f 0.1 -2653 (2.2 f 0.1) 3.1 kO.1 -33f4 (2.7 f 0.1) 3.6 f 0.1 -31+7 (3.2 f 0.1) (4.0 f 0.1) 4.3L0.1 -45 _+ 10 - 3 f 6 -19f9 -1lf17 (+ 1 f5) ( + 4 k 13) + l f l -17+7 +2&4 (+2+4) (+5f 13) -2+1 -5Of9 - 7 f 5 (+11f5) (+31 rt 15) - 7 5 4 -75+22 -27f 12 (-0.5 & 5) (-5f 13) - 3 f 5 -8Of7 -16k15 (+5+4) ( + 9 k 13) (+4+6) (+6f 16) -13+7 -91+_9 -47f20 -2Of7 - 116f 12 -69f22 ( - 1 f 5 ) (-12f18) -19+5 -113f19 -69513 (-2f 10) (- 16f28) -19f6 -157f30 -7Ok17 (-2k8) (- 20 f 23) a Organic solvent component.values are for the benzimidazole complex. Parenthesized values are for the imidazole complex; all other in the activation process. The observed non-linearity in the plot of logk;,, us. 0;' for EG-water mixtures, in contrast to eqn (3), may be attributed to (i) failure of both the point-charge model of the complex ions and the dielectric-continuum model of the reaction media and (ii) a presumably non-linear variation of the non-electrostatic component of the relative transfer free-energy term, [AG,(t.s.) - AG,(~.S.)](~~~) with the solvent composition.The values of the relative transfer free-energy term at 50 "C [calculated using eqn (3)] are presented in table 5. For both complexes it is positive, with a steadily increasing magnitude as Xorg increases. These data are consistent with the fact that the stabilizing effect of the medium on the initial state is much greater than that on the transition state, the free energy of transfer for both the transition state and the initial state being assumed to be either negative or positive. The free energy of transfer of several cations from water to MeOH-water mixtures are reported to be positive or negative, depend- ing on the nature of the cation (with respect to the charge and ligand environment) and the different approximations used.lo-12 The transfer free-energy data for the crystal violet cation [AG,(CV+) at 25 "C] reported by Kundu et a1.13 for EG-water and MeOH-water mixtures also reflect that this bulky cation with hydrophobic groups (CV' = [(CH,),NC,H,],C+} is relatively less stabilized in the former solvent.The values of the relative transfer free energy are more positive for the benzimidazole complex than for the imidazole complex at any solvent composition for the EG-water system. This leads us to believe that the solvation changes associated with both the initial state and the transition state are sensitive to the nature of the non-labile ligands imH and bzmH ; the relatively bulkier and more hydrophobic benzimidazole probably destabilizes the I, transition state in EG-water mixtures to a greater extent than imidazole (i.e.[AG,(t .s.)]+ w) (bzmH) > [AG,( t . s . ) ] ( ~ + ~ ) (imH)}. Wells'* assumes a fully82 Reactions of Coordination Complexes 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 X, Fig. 3. Dependence of the free energy of transfer [AGt(C3+)](s+w) of ~is-([Co(en),B]~+}* relative to that of cis-[Co(en),BC1I2+ at 25 "C as a function of Xorg: (1) bzmH (MeOH-water), (2) bzmH (EG-water), (3) imH (EG-water). dissociative transition state (D mechani~m)~ in order to analyse the solvent effects on the solvolytic aquation of several chloroamine cobalt(II1) complexes. The additivity principle may then be assumed to be valid for the transfer free energy of the transition (4) state to yield where [AG,(C3+)],,+,, denotes the transfer free energy of the dissociative transition state, C3+ = { cis[Co(en),(bzmH)])*. The values of the relative transfer free-energy term [AGt(C3+) - AGt(i.s.)](stw) at 25 "C were calculated from eqn (9, obtained by combining eqn (3) and (4), using the extrapolated values of the rate constants and the values of [AGt(Cl-)](s+w) reported by Abraham et al.'O (for MeOH-water, after conversion to the mole fraction scale',) and Kundu et all3 {for EG water; values of [AGt(Cl-)],s+w) on the mole fraction scale were obtained by interpolation of the data available wherever necessary).The values of [AG,(C3+) - AG,(~.S.)],,+~, were found to be negative for MeOH water at all compositions; for the EG-water system the values of this term were also negative and decreased non-linearly (see fig.3) at a higher percentage of EG. This indicates a relatively strong propensity of the tripositive cobalt(II1) species (C") to solvation, as do solvent structural changes associated with the transfer of the ionic species from the aqueous medium to the mixed solvent: ( 5 ) [AG,(t.s.)l (s+ w) = [AGt(C3+)1,s+ w) + [AGt(Cl-)I(s+ w) - AGt(i*s*)l(stw) = 2*303RT log (k,",s/k",bs) - [AGt(C1-)I(stw). The transfer free-energy calculations also show that the tripositive cobalt(II1) species (C") is more solvated in the alcohol-rich region of the mixed solvent than in water. However, this does not lead to rate enhancement, since the positive values of [AG,(Cl-)],s+,, tip the balance in favour of rate retardation with increasing XOrg.Variation of Activation Enthalpy and Entropy with Solvent Composition Fig. 4 and 5 and data in table 4 depict the variation of the activation enthalpies and entropies (AH* and AS*) with Xorg. The maxima and minima in the plots of AX* us. Xorg (X = H, S) at Xorg < 0.1 5 and the non-linear decreasing trend in these parametersA . C. Dash and N. Dash 83 I 3,4 (X = G) L 401 I 1 I I I I I 0 0.1 0.2 0.3 0.4 0.5 0.6 0.; x,,, Fig. 4. Dependence of the activation free energy (AG') and activation enthalpy (AH') of solvolysis of ci~-[Co(en),BCl]~+ (B = bzmH, imH) on the mole fraction of EG and MeOH; AG* is at 323.2 K. (1) bzmH (EG-water), (2) bzmH (MeOH-water), (3) imH (EG-water) and (4) imH. (MeOH-w ater) . beyond this mole fraction of EG and MeOH show that the solvent structural perturbations significantly modulate the values of these parameters.The values of the relative transfer thermodynamic functions [AX,(t.s.) - AXt(i.s.)](s+w) ( X = H, S) (see table 5) are also related to solvent structure. The relative transfer entropy and enthalpy at any value of X,, are sensitive to the nature of the non-labile ligand, imH and bzmH; the values of these parameters for the imidazole complex are distinctly more positive than the corresponding values for the benzimidazole complex. However, AG*(323.2 K) increases linearly with Xorg over the entire composition range; the data points for a given complex and for both solvent systems fit the same straight line reasonably well. Evidently the rate-enhancing effect due to the decreasing value of AH* is offset by the large negative value of AS*.To a large extent, the effects of AHk and AS* on the rate are mutually compensatory (see fig. 6). The large negative values of ASk being consistent with stereoretentive solv~lysis~~ corroborates the fact that the transition state is much more solvated than the initial state. The substrates under investigation possess a potentially acidic NH group (pK,, = 8.6 and ca. 10 for the coordinated benzimidazole and imidazole, respectively3? 16) adjacent to the reaction site (i.e. the Co-Cl dipole), and since the methyl group and the bulky benzimidazole species can take part in hydrophobic ass~ciation'~ we believe that solvation in the initial and transition states will involve cooperative hydrogen bonding in which the solvent molecules, Co-C1 and N-H dipoles of the imidazole or benzimidazole ligands take part as depicted in fig.7. [Note that the model of front-side solvation was postulated by Adamson and Basolo18 in the context of the aquation of (NH3),CoC12+.] This effect is likely to be strengthened at the NH site in the transition state, since development of positive charge at the cobalt(r1r) centre will accentuate the dipole-dipole interaction. A similar situation involving EG84 Reactions of Coordination Complexes 0 1 . I d I - ; ;c " -100- * \ 5 -200 - I I I I I 1 lI J 0 0.1 0.2 0.3 0.4 0.S 0.6 0.7 Fig. 5. Dependence of the activation entropy of the solvolysis of ~is-[Co(en),BCl]~+ (B = bzmH, imH) on the mole fraction of EG and MeOH. (1) bzmH (EG-water), (2) bzmH (MeOH-water) Xorg -180 -160 -140 -120 -100 -80 -60 -40 -20 0 M*/J K-' mol-' Fig.6. Plot of AH* kJ mol-l against AS* J K-' mol-l: 0, bzmH (MeOH-water); A, bzmH (EG-water) ; 0, imH (EG-water). molecules probably prevails in EG-water mixtures, and presumably the initial and transition states are preferentially solvated by EG, since it can effectively participate in hydrogen bonding with the NH and Co-Cl dipoles by utilizing both of its OH groups. In this context one should note the relatively more positive values of AS* (ASiGwater > AS&eOH-water, see table 4) at all compositions of EG-water for the benzimidazole complex. Furthermore, the large decrease of AH* with increasing XMeOH for cis-[Co(en),- (bzmH)C1I2+ (see fig. 4 and table 4) reflects the energetic role of the solvent cosphere ofA.C. Dash and N . Dash 85 cis Fig. 7. Front-side initial state solvation in which SH is MeOH, H,O or EG. this complex in indirectly influencing the development of the highly polar transition state. The overall values of AH* and AS* may be taken to be composites of the reaction component and the solvational component1' (AX&verall = AX; + AX,'). The validity of the compensation law (see fig. 6 ) is therefore in keeping with the fact that perturbations of the reaction zone and the solvent network remain proportional to each other with increasing Xorg, so that the isodelphic and the lyodelphic components of AH* and AS* correlate well with each other over the entire composition range of both the solvent systems used.20 This work was supported by a grant from CSIR to A.C.D.N.D. thanks the CSIR for a Junior Fellowship. References 1 2 3 4 5 6 7 8 9 10 1 1 12 13 14 15 16 17 18 19 20 Part 1 : A. C. Dash and N. Dash, J. Chem. Soc., Faraday Trans. 1, 1987, 83, 2505. (a) E. Grunwald and S. Winstein, J. Am. Chem. SOC., 1948,70,846; (b) T. W. Bentley and G. E. Carter, J. Am. Chem. SOC., 1982, 104, 5741. A. C. Dash and S. K. Mohapatra, J. Chem. Soc., Dalton Trans., 1977, 1207. I. J. Kindred and D. A. House, Inorg. Chim. Acta, 1975, 14, 185. G. Akerlof, J. Am. Chem. SOC., 1932, 54,4125. J. Burgess and M. G. Price, J. Chem. Soc. A, 1971, 3108 and references therein; G. Thomas and L. A. P. Kane-Magurie, J. Chem. Soc., Dalton Trans., 1974, 1688. B. G. Cox, G. R. Hedwig, A. J. Parker and D. W. Watts, Austr. J. Chem., 1974,27, 477; A. J. Parker, Electrochim. Acta, 1976, 21, 671. C. G. Swain, M. S. Swain, A. L. Powell and S. Alunni, J. Am. Chem. Soc., 1983, 105, 502. C. H. Langford and H. B. Gray, Ligand Substitution Processes (W. A. Benjamin, New York, 1965). M. H. Abraham, T. Hill, H. C. Ling, R. A. Schulz and R. A. Watt, J. Chem. SOC., Faraday Trans. 1, 1984, 80, 489. C. F. Wells, J. Chem. Soc., Faraday Trans. 1, 1973, 69, 984. M. J. Blandamer, J. Burgess, B. Clark, P. P. Duce, A. W. Hakin, N. Gosal, S. Radulovic, P. Guar- dado, F. Sanchez, C. C. Hubbard, Ezz-E. A. Abu-Gharib, J. Chem. Soc., Faraday Trans. 1 , 1986, 82, 1471. U. Mendal, S. Sen, K. Das and K. K. Kundu, Can. J. Chem., 1986, 64, 300; A. K. Das and K. K. Kundu, Ind. J. Chem., Sect. A , 1978, 16, 467. A. E. Eid and C. F. Wells, J. Chem. SOC., Faraday Trans. 1, 1985, 81, 1401. M. L. Tobe, Znorg. Chem., 1968, 7 , 1260. A. C. Dash, M. S. Dash and S. K. Mohapatra, J. Chem. Res., 1979, ( S ) 354; (M) 4531. F. Franks, Water-A Comprehensive Treatise, ed. F. Franks (Plenum Press, New York, 1973), vol. 4, A. W. Adamson and F. Basolo, Acta Chem. Scand., 1955, 9, 1261. M. J. Blandamer, Adv. Phys. Org. Chem., 1977, 14, 247. E. Grunwald, J. Am. Chem. Soc., 1984, 106, 5414. p. 1. Paper 612340; Received 3rd December, 1986

ISSN:0300-9599    

DOI:10.1039/F19888400075

出版商:RSC    

年代:1988

数据来源: RSC