摘要:

J.C.S. Faraday I, 1980,76,2552-2557Electrolyte Diffusion at Very Low Concentrations inIonized WaterBY PENTTI PASSINIEMI," SIMO LIUKKONEN AND ZOLTAN NOSZTICZIUS-/-Helsinki University of Technology, Department of Chemistry,SF-02150 Espoo 15, FinlandReceived 6th November, 1979The diffusional behaviour of an electrolyte Mv +Xv- inwater is studied up to concentrations wherethe effect of the ionization must be taken into account. The mutual diffusion in this system changesdue to the dissociation of water under the decreasing electrolyte concentration, first to multi-component diffusion and then, at concentrations < lo-* mol dm-3, to tracer diffusions of the electro-lyte ions.By applying Nernst-Planck equations the approximate values of the proper difiusion coefficientsare calculated.These rather simple calculations confirm the phenomena which were observedexperimentally by Mills at low concentrations in the system H20+ MgBrz at 298.15 K.Very little attention has been paid to transport phenomena in very dilute aqueoussolutions of electrolytes where the ionization of water can no longer be neglected.Measurements in these solutions are often very difficult and, from the theoreticalpoint of view, the reference velocity may give rise to questions.However, there is one very interesting series of measurements from 1962, in whichthe diffusion has been studied up to a concentration of 3 x 10-8moldm-3 usingradioactive 82Br in the analysis. In these experiments of Mills,l magnesium bromidediffused out of a capillary into pure water.The results of this series were analysedtheoretically by Woolf.2As the treatment of the transport processes in very dilute aqueous electrolytesolutions raises some fundamental questions we want to consider this problem anew.Qualitatively the diffusion in such a system is clear as has already been discussed byMills' and Woolf et aL2* In concentrated binary solutions a mutual diffusionprevails but, on the other hand, in very dilute solutions (< mol dm-3) thereexist one or more tracer diffusions. Between these extreme cases or in the inter-mediate region the concentrations of H+ and OH- ions due to the water are com-parable with the concentration of the electrolyte. Then there is essentially a multi-component system and coupled diffusions appear.To get a closer and more quantitative description of these diffusion phenomenawe will apply Nernst-Planck equations to the ionic flows.The component flowswith proper diffusion coefficients are then estimated in terms of the ionic molar con-ductivities at infinite dilution.THEORYThe isothermal-isobaric system to be considered is composed of asociated electrolyte Mv,Xv- in partially ionized water. The electrolytest Present address : Technical University of Budapest, Budapest, Hungary.2552fully dis-dissociaeP. PASSINIEMI, S . LIUKKONEN AND 2. NOSZTICZIUS 2553according to eqn (1) and (2) :M,+X,- + v+MZ+ +v-X"-- (1)H 2 0 + H++OH-. (2)Kw = C H C o H = mo12 dm-6. (3)The ionic product of water is at 25°CFor simplicity, the symbol for an ionic species is written without charge, e.g., Hinstead of H+, etc.In studying the diffusion of the electrolytes the approximations of Nernst-Planckequations are applied to the ionic flows in the concentration range to be discussed.The flows are 4,-ji = Di(VCi +ziFCiV4/RT) i = H, OH, M, X,where j i , Di, Ci and zi are, respectively, the flow, the diffusion coefficient, the con-centration and the charge of the ionic species i.4 is the so-called diffusion potential.Di is defined by(4)Di = RTA?/(z?P), ( 5 )in which A? is the ionic molar conductivity at infinite dilution. We assume that theseconstant A? values can also be applied to ionic conductivities at very dilute con-centrations.In order to describe the diffusion in the most general case we start with the ternarysysteminstead of the binary one defined by eqn (1) and (2).* In Nernst-Planck equationsof this system V4, VC, and VC,,, respectively, are eliminated by the conditions ofthe zero electric current densityH2O + H,z-,X(l) + M(OH)z+(2) (6)X _ _ziFji = 0, (7)i = Hthe constancy of the ionic product of water [eqn (3)] in the formand the electroneutrality in the formX c zivc, = 0, (9)i = HIn eqn (7) and (9) the sum includes all the ions, i.e., H, OH, M and X.The ionic flows are thusX-jk = D, C (~ik-(ZiZkCk/X)[Di-(KDOH+ ~ , ) / ( l +K)I}VCi k = M, X, (10)(11)i = MX-j, = [D,Ai/(1 +K)I 1 Zi(1 -(CHIX)[(KDO,+ ~ d - ( 1 +K)DiI}VCii = MI = H,OH and AH = - 1 , AOH = K.* It is necessary to make clear the difference between the total concentration of the hydrogen ionsand the concentration C1 of the component Hlz-1X [in eqn (6)].It means in this treatment thatCH # 12-1 CI. The same is also true for the hydroxyl ions2554 DIFFUSION AT VERY LOW CONCENTRATIONSThe sum in eqn (10) and (11) is taken only over the ions M and X. &k in eqn (lo)is a Kronecker delta and K and x in eqn (10) and (1 1) are, respectively, given byIf we choose for the independent flows jM and j , we can interpret the coefficientsof VCM and VC, in eqn (10) to be the four diffusion coefficients of the ternary system(6). These diffusion coefficients are concentration dependent and they include theeffect of the dissociation of water. Thus the flows areMUTUAL DIFFUSION CM, C, CH, COH ; Z+CM = -2-c~With these conditions we have&/v+ =jx/v- = JMy+xv- ; j H = j o H =I 0,and the approximate mutual diffusion coefficientequals the limiting value of N e r n ~ t .~DMly+Xv- = [(.+ -z-)DMDXl/(z+DM-z-DX)TRACER DIFFUSION CM, Cx < CH, C ~ HIn this case we see from eqn (14) thatThere are two independent tracer flowsDkf = 6 k i D k i, k = M, x. (1 7)(184( W- j M = -JM(OH)z+ = DMvcM.-jx = -JH,.-IX = D X V C Xin which the tracer diffusion coefficients DM and Dx are given in eqn (5).case both j , and joH approach zero.In thisINTERMEDIATE REGION CM and Cx comparable to CH and C,,We now have ternary diffusion in system (6) with non-zero cross-diffusion coefficients.There are also non-zero concentration gradients for H and OH ions (pH is notspatially constant).THE SYSTEM H20+MgBr2We have chosen this system as an example to which we will apply the abovetreatment.The system is the only one for which there are experimental data fromthe binary (mutual) diffusion to the tracer diffusion. In calculations the followingvalues were used for @' :6 A"(H+) = 349.81 ; A"(0H-) = 197.8; A"(Mg2+) =106.10; j103(Br-) = 78.14 cm2 Q-' mol-'P . PASSINIEMI, S . LIUKKONEN AND Z . NOSZTICZIUS 2555The diffusion coefficients thus obtained are given in table 1. They have beencalculated in the whole concentration region at the point which corresponds topH = 7.00. C in table 1 is the concentration of MgBr, at this point.TABLE 1 .-ESTIMATED DIFFUSION COEFFICIENTS IN THE SYSTEM H20 + MgBrz AT 298.15 KC" Dllb D12b D21b D22b, DBrb DMgb DBrMg?10-310-410-510-710-91 o-6lo-*10-l05.1885.1825.1254.6523.0872.2232.0962.082- 7.852 - 1.054- 7.837 - 1.052- 7.693 - 1.033- 6.496 - 0.872-2.542 -0.341- 0.359 - 0.048- 0.037 - 0.005- 0.003 - 0.0013.370 1.2623.365 1.2623.3162.9101.5690.8280.719 2.077 0.7090.706 2.082 0.706~~ ~ ~~~~~~a Units : mol drr3 ; units : lop9 m2 s-l.The mutual diffusion coefficient DMgBr2 is determined from eqn (16) but it can alsobe obtained by combining Dki, under the assumptions of the mutual diffusion caseabove, as follows :DMgBrz = *(2Dll + O12) = (2D21 + D22)- (19)In eqn (19) subscript 1 denotes, according to system (6), the component HBr and 2the component Mg(OH),. At concentrations z mol dm-3, DI2 and DZl arenearly zero.If we assume that the concentration gradients VC, and VC,, are alsozero we can combine Dki to find the approximate values for the tracer diffusioncoefficients DMg and D,, as follows :These values are given in table 1.The mutual diffusion coefficient DMgBr2 can be determined by labelling, e. g.,Br ion (or Mg ion) in MgBr, and by applying the diaphragm-cell or open-endedcapillary method. The measurements for the mutual diffusion coefficient can bemade as given in table 1 up to concentrations z mol dm-3 and from these resultsthe limiting value of Nernst can be reached with the usual extrapolation.On the other hand, the limiting values of the tracer diffusion coefficients DMg(Mg Iabelled in MgBr,) and DBr (Br labelled in MgBr,) are obtained with two differentmethods.Firstly, under the conditions of the usual tracer- (or better the self-)diffusion experiment where the total concentration of the electrolyte is constant.The limiting values can almost be found at concentration z mol dm-3. Thedissociation of water has no effect in this case. Secondly, the same limiting valueDBr is achieved by measuring the diffusion of Br-labelled MgBr, or of any otherelectrolyte containing Br ion into pure water at concentrations < mol dm-3.Correspondingly, D,, is obtained with Mg-labelled MgBr, or with any other electro-lyte containing Mg2+-ion.The four diffusion coefficients in the intermediate region are also, in principle,measurable with the open-ended capillary method by applying, e.g., the ideas ofToukubo and Nakanishi.'DBr *(2D11 +D12) ; DMg z 2D21i-0222556 DIFFUSION AT VERY LOW CONCENTRATIONSDISCUSSIONTo complete our analysis of the diffusion in the system H20 + MgBr, we have alsocalculated the diffusion coefficients in the systems H20 + HBr and H 2 0 + Mg(OH),.In both cases the treatment above shows that there is only one independent flow in thewhole concentration region.In the former system the flow of Br- is given by-jBr = -JHBr = DBr(l + ( c H B r / x > ( l + K)-l[(KDO€l + DH) -+K)DBrl)VCHBr (20)and in the latter system the flow of Mg2+ is given by-jMg = -JMg(OH)2 = +(4cMg(OH,,/x)(1 + K)-l+ DH) - (l + K)DMg]}vcM~(OH)2- (21)In eqn (20) and (21) we have only one measurable diffusion coefficient in the wholeconcentration region.These coefficients could be determined, e.g.? with the open-ended capillary method by using the labelled Br- (or Mg2+) ion for analysing thechanges in the electrolyte concentration. At higher concentrations (up to zi moldm-3) the value measured approaches the limiting value of the mutual diffusioncoefficient. At concentrations under mol dm-3 this mutual diffusion coefficient,i 3.0 iI1 .o]-10 - 8 - 6 -4log (C/mol dm-3)FIG. 1.-Calculated diffusion coefficients at very low concentrations. The full lines (a), (c) and ( d )are, respectively, the Onsager limiting laws of the aqueous binary diffusion for HBr, Mg(OH), andMgBr,. The full lines (6) and (e) are the Onsager limiting laws for the tracer diffusion of Br- andMg2+ in aqueous MgBr2.The broken lines (f) and (9) describe the change of the binary diffusioncoefficient in aqueous HBr and Mg(OH)2, respectively. The dotted lines (h) and (i) give the sums(Dll+0.5 D12) and (2021+022) in the system H20+MgBr2 with respect to the concentration atpH = 7.00 [see also the text after eqn (19)]P . PASSINIEMI, S . LIUKKONEN AND Z . NOSZTICZIUS 2557however, ought to change quite rapidly according to eqn (20) and (21) and in a solu-tion concentration z mol dm-3 the values of the tracer diffusion coefficients(DBr and DMg) could be reached. This behaviour can be seen in fig. 1 where thediffusion coefficients in the system H,O+MgBr, are also shown.There one seeshow MgBr, is divided into HBr and Mg(OH), as its concentration decreases. Notethat the traditional extrapolation in mutual diffusion coefficients is made from experi-mental results in the concentration range of z 10-3-5 x mol dm-3 to zeroelectrolyte concentration. The limits thus obtained do not contain the effect dueto the water dissociation.Eventhe small maximum in the diffusion coefficient could be explained by the pH-changedue to the possible adsorption of Mg-ions on plexiglass. BeneS and KopiEka *have shown that especially divalent cations are adsorbed readily on plexiglass invery dilute solutions. Unfortunately, Mills measured only the diffusion in the systemH,O + MgBr, where the interpretation of his diffusion coefficients in the region105-10-8 mol dm-3 is not unambiguous [cf.eqn (10) above].If we compare the diffusion coefficients given in our table 1 with those in table ofWoolf’s paper there are quite large deviations. Woolf used the method of Wendtfor the approximate treatment in system (6). However, in principle that method isanalogous with the Nernst-Planck equations. O In Woolf ’s paper the treatment ofdiffusion is divided into three regions. It seems to us that the connection of themutual and the tracer diffusion regions with the intermediate region is not clear.In the present paper the gradual change from the mutual to the ternary diffusionand further to the tracer diffusion is obvious.We want to emphasize that the Nernst-Planck equations are also written here tohydrogen- and hydroxyl-ions.This means the electric conductivity of the systemalso includes the part due to the dissociated water as given in eqn (12b). Therefore,if the concentration of the electrolyte MX goes to zero the transport numbers ofMZ+- and Xz--ions also go to zero. This corresponds to the result in molten saltsbut now the non-measurable transport numbers of H+ and OH- in pure water donot equal one and zero but are related to their ionic conductivities.To test the above derivations for the diffusion coefficients one would need diffusionmeasurements in very dilute solutions especially in systems of the form H20 + H,,-,Xand H 2 0 + M(OH),+.Our calculations confirm the main points of Mills’ experimental data.lWe acknowledge scholarships made available by the cooperation between theP. P. thanks the Finnish Cultural Technical University of Budapest and Helsinki.Foundation and S . L. the Finnish Academy for the scholarships.R. Mills, J. Phys. Chem., 1962, 66, 2716.L. A. Woolf, J. Phys. Chem., 1972,76, 1166.L. A. Woolf, D. G. Miller and L. J. Gosting, J. Amer. Chem. SOC., 1962, 84, 317.W. Nernst, 2. phys. Chem., 1888, 2, 613 ; 1889, 4, 129.M. Planck, Ann. phys., 1890, 40, 561.R. A. Robinson and R. H. Stokes, Electrolyte SoZutions (Butterworths, London, 2nd edn,1959), p. 463. ’ K. Toukubo and K. Nakanishi, J. Phys. Chem., 1974,78,2281.* P. Bene’s and K. KopiEka, J. Inorg. Nuclear Chem., 1976, 38, 2043.R. P. Wendt, J . Phys. Chem., 1965, 69, 1227.lo H.-J. Schonert, 2. phys. Chem. (Frankfurt), 1967, 54, 245.l 1 C. Sinistri, J. Phys. Chem., 1962, 66, 1600.(PAPER 9/1776

ISSN:0300-9599    

DOI:10.1039/F19807602552

出版商:RSC    

年代:1980

数据来源: RSC

摘要:

J.C.S. Faruday I , 1980,76,2558-2574Infrared Investigation of Ionic Hydration inTon-exchange MembranesPart 1 .-Alkaline Salts of Grafted Polystyrene SulphonicAcid MembranesBY LEON Y. LEVY, ANDR~ JENARD AND HENRI D. HURWITZ"Laboratory of Electrochemical Thermodynamics, Faculty of Science,Universitk Libre, Brussels, BelgiumReceived 12th November, 1979Infrared spectroscopic measurements have been performed on thin ion-exchange membranes whichconsist of alkaline salts of polystyrene sulphonic acid grafted on a Teflon FEP matrix. The membranesunder investigation were placed in isopiestic equilibrium with water vapour in the sampling cell. Thewater content of these FEP-PSSA systems has been determined at 25°C by measuring the integratedabsorbance of the sorbed-water bending vibration as a function of the nature of the cation.The depen-dence of the position of the water stretching and bending vibration bands and of the symmetricvibration of the SO, groups with the nature of the cation and the water content has led to variousconclusions concerning the ionic interaction and hydration in the membrane. The peculiar behaviour ofLi' is due to its ability to interact strongly with the anion. A model of ion clustering has been proposedto explain the spectra of water in the presence of large alkali ions.Some infrared spectroscopic measurements on polystyrene sulphonic acid mem-branes were carried out by Zundel in 1969' in order to investigate the ion-solventinteraction in the polymer. Since then, few i.r.spectroscopic studies have beendevoted to the elucidation of the effect of water and of the exchangeable ion inion-containing polymeric membranes. This is surprising as a vast literature existsand the industrial and analytical applications of such ion-exchange materials haveincreased considerably. Among the polymeric ion-exchange membranes which areof special practical use to date, one finds films containing polystyrene sulphonicacid (PSSA) grafted onto fluorocarbon backbones? and so-called perfluorosul-phonic acid membranes made out of saponified copolymers of sulphonyl fluoridevinyl ether and tetrafluoroethylene.1 In our spectroscopic investigation, we havefocussed our attention on this type of material. The present publication describesthe i.r.spectroscopic analysis performed on the alkaline salts of PSSA grafted on aTeflon FEP matrix. The future publications will present spectroscopic measure-ments made on other salts of PSSA and on salts of perfluorosulphonic acid(Nafion) membranes.It is known from the investigations of Zunde12 and from several contributionsdealing with strong electrolytes that the vibrational spectra of H 2 0 3 and of SO:-may be deeply perturbed under the influence of the polarizing field produced by anearby cation. For instance, the OH stretching band and HOH scissoring band ofwater will both be shifted to different wavenumbers with ditferent magnitudes. Thusit is worthwhile to corfelate such effects with a model of ionic solvation andt AMF C322 PSSA grafted on polytrifluorochloroethylene (Kel F) or on fluorinated ethylene propy-$ Nafion (registered trade mark) produced by Dupont.lene copolymer (FEP).255FIG.l - - - ( ~ i ) : Electron micrograph of a FEP film surface ( x 15600). (h) Electron micrograph of the surfaceof a FEP membrane grafted with PSSA in uranyl salt form ( x 10300).To face page 2559L. Y . LEVY, A. JENARD AND H . D . HURWITZ 2559interaction. In the case of ion-exchange membranes, such an approach shouldprove particularly rewarding if the microscopic interpretation is successful inexplaining such macroscopic properties as equilibrium swelling, wetting and ionicselectivity. In this respect, the interpretation of the value of the ionic separationfactor of ion-exchange resins and membranes relies nowadays on various theories,some based on more or less crude statistical mechanical others onpurely thermodynamic deduction^.^ The complexity of combining several types ofionic interactions which are competing with hydration in the ionomeric materialalso produces models which are frequently restricted to qualitative and empiricaltheories by which the ability of the model to yield consistent counterion sequencesis emphasized more than its agreement with any statistical thermodynamic ormolecular computation.OSpectroscopic arguments might help to reconcile the various treatments. Conse-quently, we endeavour here to substantiate phenomenological properties (such asequilibrium wetting and ionic selectivity) in terms of molecular models of ionicsolvation and interaction in the polymeric membrane.Also, the specific influenceexerted on these effects by the inert macromolecular matrix will be explored.EXPERIMENTALPREPARATION OF SAMPLESMembranes of 62 f 1 pm thickness of PSS grafted on FEP" were used. The thickness of asample of non-grafted film was determined from its interference pattern in the i.r. spectrum. Itwas 50 pm thick within f 3%. All i.r. spectroscopic measurements were carried out on samplesof 32 and 15 mm radius.The membranes, initially in their acid form, were placed in 1 mol dmP3 solutions of alkalichloride for at least three periods of 4 h. All solutions were prepared with Merck pro anulysisalts. After being removed from the solution, the membranes were washed with triply distilledwater and carefully wiped with filter paper.DETER MI N ,\ T I O N OF EX C H A N GE C A P A C I T I E SThe exchange capacity of the samples prepared for i.r.spectroscopic studies were obtained bycoulometiic microtitration carried out following the method described by Sansoni.' A Ptelectrode of 1 cm2 area was used as hydrogen ion source. The cell assembly also contained acombined Ingold microelectrode (type HA 405 MJNS) and a Pt counterelectrode in a separatecompartment connected with the main compartment by means of a saturated K2S04 agar-agarbridge. A potentiostat (Tacwsel PRT 3001) was used as a galvanostatic source providingcurrent intensities ranging from 500 to 1000 pA with a precision better than 1 pA. The pH inthe cell was kept constant.The average value of the exchange capacity was 1.15 meq g- andthe scattering of the results was ~ 2 . 5 % for the various samples irrespective of the salt, whichindicates a small degree of inhomogeneity of the membrane. Such inhomogeneity, due to itsgrafting on the FEP film, could be studied by electronic microscopy. The membrane surfacewas observed using the technique of carbon replica shaded with silver at an angle of 15". Anelectron micrograph for an ungrafted film was taken in the same condition. Significant areas ofthe surface of an ungrafted film [fig. l(a)] and grafted film [fig. l(b)] are shown. By inspectionof the first micrograph, one: can exclude a microphase separation. Homogeneous distribution ofinterwoven chains is observed.The micrograph of the grafted membrane reveals the existenceof slightly swollen domains heterogeneously distributed on the surface. The diameter of these* Progil, France; ref. (2.50-7-702560 IONIC HYDRATION I N ION-EXCHANGE MEMBRANESdomains is up to 20.3 pm. These domains are presumably regions of high PSSA penetrationsurrounded by fluorinated backbones presenting a much lower graft density. The techniqueused did not detect any supermolecular structure as found by Ceynowa12 on a PTFE film.DRYING OF MEMBRANESThe membrane was dried in two successive operations. First, the membrane has been placedin a vacuum Torr) in the presence of P205 at 80‘C and, secondly, the dry membrane wasexposed to i.r. light at a temperature of 2 50‘C in the spectrometer sampling cell through whicha flow of dry air ( < 4 p.p.m.of H 2 0 at 25‘C) was delivered. The passage of dry air continueduntil no change was observed for a period of 24h in the continuously recorded i.r. waterabsorption band of the membrane. Such conditions of dryness were achieved after 4-6 daysdepending on the nature of the exchangeable cation. The final spectrum of the so-called“thoroughly dried membrane” indicated that there was still some residual hydration whichcould not be removed even after two or three weeks of a similar repetitive drying operations.DbSCRIPTlON OF THE I . R . SAMPLING CELLA sampling cell was designed to meet the following requirements: (i) the membrane underinvestigation must be placed in isopiestic equilibrium conditions in the cell; (ii) the cell mustdispose of a large thermal capacity and must be thermally well isolated; (iii) the saturatorscontaining solutions providing the selected water vapour pressures must be easily interchange-able.The body of the cell, as sketched in fig.2, is made of brass. It has cylindrical symmetry and iscomposed of two pieces (1 and 2) which can be combined by means of a ring screw and twoO-rings which insure a good vacuum seal. Inside the cell is placed a membrane holder of ironcovered with cadmium (3). The membrane (4) is fixed on its holder by means of a magneticribbon (5). Two nozzles (6) permit the circulation of thermostatic Auid through the concentricjacket of the cell. The inside of the cell is joined to a saturator and to the vacuum pumpthrough a nozzle (7). The two AgCl windows (8) have a PTFE frame (9) which insulates themFIG.2.--Exploded view of the i.r. sampling cellL . Y . LEVY, A . JENARD A N D H . D . HURWITZ 256 1from the metal body of the cell. This frame is pressed on the cell by means of a Teflon flatO-ring screw (10 and 11). The path length from window to window is 5 cm. The saturatorconsists of simple Pyrex tube sealed on the entry of a three-way stopcock connected to thepump and to nozzle (7).DETERMINATION OF WATER CONTENTHydrated membranes were titrated by the Karl Fischer method. For this purpose sn auto-matic K.F.4 Beckman apparatus was modified slightly in order to satisfy the anhydrous con-ditions for the titration agent and the titration cell during the experiment.The membranes aresubjected to a fast spontaneous rehydration or converse slow dehydration whenever theselected isopiestic water absorption equilibrium conditions are enforced, which happens duringthe transport of the membrane to the Karl Fischer cell. These processes prohibit accuratedetermination of the degree of wetness by the Karl Fischer titration unless the membrane istaken at equilibrium exchange with the laboratory atmosphere at 25°C. For this reason, it wasnecessary to conceive an indirect way of water content determination in membranes which areplaced at equilibrium at a relative humidity different from the laboratory atmosphere. Thismethod will be described later in the results section.SP E C TR 0 S C OP 1 C ME A SUR EME N T S U N D ER I SOP I E S T I C CON U 1 T I ON SSamples previously tested, under identical conditions, for their water content and exchangecapacity were thoroughly dried in the i.r.sampling cell. A progressive rehydration was carriedout stepwise by placing saturated electrolytic aqueous solutions of increasing water vapourpressure in the saturator. It is assumed that the water absorption equilibrium is reached whenthe water i.r. absorption bands remain unchanged for 24 h after connecting the saturator andthe cell, both at 25°C. The spectra of the membranes were recorded with a Beckman i.r. 9double-beam spectrophotometer. An identical cell, without the membrane but set at the samevapour pressure, was inserted in the reference beam.In order to obtain good accuracy in bandfrequency determination, a slow scanning of 20 cm-' min-' was used and the bands wereenlarged using an amplification factor of 4 or, in some cases (OH stretching absorption band), afactor of 2. This extends the accuracy to z 2 I cm-' for the narrow bands and yieldsin the most unfavourable cases (the very large water stretching bands or shoulders). It shouldbe stressed that for the hydrated membranes at equilibrium with the laboratory atmosphere at2 5 T , only the first of the recorded spectra was retained since a slow process of dehydrationcould occur under the influence of the i.r. radiation.Each complete set of investigations, as a function of the membrane wetting and the nature ofthe countercation, was made on the same sample in order to avoid the effect of inhomogeneitiesin the film.The accuracy of frequency measurements has, however, been checked by repeateddetermination on several samples cut from different films.5 cm-ASSIGNMENT OF THE 1 . R . ABSORPTION BANDSIn fig. 3(u) is shown a selected spectrum of the K + salt of a FEP-PSSA film takenat equilibrium with an atmosphere of relative humidity p / p o = 0.98 at 25'C. Thespectrum of the FEP film taken at the same equilibrium conditions is given forcomparison in fig. 3(h). The assignment of the most important spectral bands are asfollows :The broad band in the region 3700-3300cm-1 is ascribed to the stretchingvibrations of H20. In the case of the ion-exchange membrane, the position of themaximum of absorption and the intensity of this band depend on the nature of thecation as well as the degree of hydration.Thus this band can be attributed tomolecules involved in ionic solvation inside the membrane. The presence of freeOH groups, thus not involved in a hydrogen bond, might produce some shouldersat z3600cm-'2562 IONIC HYDRATION I N ION-EXCHANGE MEMBRANES1 I I I I I lOOr I I I II I , I 14000 3600 2 800 2000 1600 1200 800 400wavenumber/cm -FIG. 3.--(a) 1.r. spectrum of a K' salt of the FEP-PSSA membrane at 98:< relative atmospherichumidity. (b) 1.r. spectrum of a FEP film.The bands at 3060 and 3025 cm-' are caused by the stretching vibrations of the>CH groups in the benzene ring.These bands cannot be related to the individualvibrations of these groups since they are coupled with other vibrations appearing inthe system.I3The bands at 2924 and 2851 cm-' are assigned, respectively, to the asymmetric andsymmetric stretching vibrations of the -CH2- group. l4The band at 2398 cm- ' is characteristic of the first harmonic of the CF2 stretchingvibration. The shape and intensity of this band depends on the type of ion and thewater content in the membrane. It has been suggested that the thickness andcrystallization of the Teflon film may be determined by means of this band.15-17 Inthe case of a grafted film, this band might, ,however, hide a weak absorption banddue to the sorbed hydrogen-bonded water molecules.The four weak bands at 1938, 1845, 1792 and 1725 cm-' (mixed with an interfer-ence pattern) are due to combination vibrations and overtones of the out-of-planebending vibration of the > CH group of the benzene ring.'The band at 1640 cm- ' corresponds to the scissor vibration of water. By analysis ofcationic influences on this band, specific characteristics of ionic hydration phenom-ena in the membrane are detected.Bands at 1599, 1494 and 1411 cm-' are caused by the skeletal stretching vi-brational of the benzene The band at 1445 cm-' is the scissor vibration ofthe -CH2- group.14The very intense band ranging from 1100 to 1300 cm- ' corresponds to the stretch-ing vibration of the CF2 group.This band is masking the antisymmetric stretchingof the SO, ion (thoroughly analysed by Zundel').The band at 1040cm-' is attributed to the symmetric stretching vibration of theSO, ion in mesomeric form of C30 symmetry.The position of this band isinfluenced by the cation.The band at 1011 cm-' is ascribed to the ion plane bending vibration of the >CHgroup of benzene." This peak is shifted towards 1000 cm-' as a function of thedegree of dryness of the membrane.The band at 981 cm-' is characteristic of Teflon FEP and might be due to the-CF3- group asymmetric stretching vibration.The band at 900 cm-' is due to the stretching vibration of the single S-0 bondappearing in the SO, non-mesomeric form or HS03 acidic form. It is observedwith alkaline salts but its very weak intensity does not allow any significant inter-pretation.2L .Y . LEVY, A. JENARD AND H . D. HURWITZ 2563The band at 831 cm-l is due to the out-of-plane bending vibrations of the twopairs of > CH groups on each side of the para-disubstituted benzene ring. l9The assignment of the peak at 773 cm-' is still unsettled. By comparison with thespectra of ortho-toluene sulphonic acid, Zundel has suggested that this band corre-sponds to the out-of-plane deformation of four CH groups in the benzene skele-ton.' Note that this band is found at the same location as an absorption peak ofTeflon FEP.RESULTSINFLUENCE OF WATER CONTENT ON THE INTENSITY OF THE WATER I.R.ABSORPTION BANDSIn the case of hydrated membranes taken at equilibrium with the laboratoryatmosphere, the maximum absorbances and integrated absorbances of the waterstretching and bending vibration bands have been plotted in fig.4 and 5 as afunction of the amount of water measured by the Karl Fisher titration method. Acomputer program has been used to integrate the absorbances as a function ofnH2OFIG. 4.-OH stretching vibration integrated and maximum absorbance of the hydration water as afunction of the degree of hydration of the membrane in its alkaline and alkaline-earth salts. 0, Li; @, Na;0, K ; U, Cs; @, Mg; 8, Ca; +, Sr; 0, Ba2564 I O N I C HYDRATION I N ION-EXCHANGE MEMBRANESlTH 0FIG. 5.-OH bending vibration integrated and maximum absorbance of the hydration water as afunction of the degree of hydration of the membrane in its alkaline and alkaline-earth salts.Symbols asin fig. 4.wavenumber. This computation is made by subtracting from the integrated valuesthe contributions of, respectively, the shoulder at 3025 cm-I and the bands at 3060,2924, 2851 and 1600cm-'. The optical integrated densities due to these variouspeaks are calculated from the spectrum of the Cs+ salt of the thoroughly driedFEP-PSSA membrane, for which a minimum of wetting is observed. The deconvo-lution in this system is based on the assumption of a gaussian shape for thestretching and scissoring vibration bands of water. As a result of these calculations,we have ascertained that a linear relationship exists between the integrated absorb-ance of the bending vibration mode Ad and the amount of absorbed water, irrespec-tive of the nature of the cations.An equivalent simple behaviour is exhibited neitherby the integrated absorbance of the stretching vibration mode A , , which shows twodifferent linear segments depending on the charge of the countercation, nor by themaximum absorbances of both modes.21Using these results we have computed the relative water content of any samplesubjected to various hydration conditions, by measuring its integrated absorbanceAh and assigning a relative value of 100 to the A6 value of the membrane withME2+ salt taken at its maximum degree of swelling. The results, presented in theform of absorption isotherms, are depicted in fig. 6 for alkaline saltsL . Y . LEVY, A . JENARIJ AND H . D. HURWITZ 2565relative humidity (x)FIG. 6.- Absorption isotherms at 25'C of the Li' (0) and K' (U) salts of the FEP-PSSA membrane.DEPENDENCE OF BAND POSITIONS ON THE NATURE OF THECATION AND ON THE WATER CONTENTPositions of the maxima of the intense bands corresponding, respectively, to thesymmetric vibration v, (SO;) of the SO, group, the stretching vibration vOH andbending vibration aOH of the hydration water molecules in FEP-PSS membranes at7% relative humidity and 25°C are recorded in table 1.With our results we alsoshow some group vibrations obtained by Zundel' for salts of ungrafted PSS mem-branes. The shift in frequency of the maxima of the absorption bands recorded as afunction of the water content for the various alkaline salts of FEP-PSS membranesis given in fig. 7-9.As regards the effect of water content on other vibrational group absorbances, letus note that some influence has been detected on a shoulder which exists to thelow-frequency side of the CH bending vibration of the benzene ring atw 1000 cm-'.DISCUSSIONThe FEP films possess a high mechanical resistance and are extremely hydro-phobic. Hence we do not find any trace of hydration, as revealed by i.r.waterabsorption bands, in pure FEP films kept for three weeks in contact with anTABLE 1 .-VIBRATION FREQUENCIES FOR POLYSTYRENE SULPHONIC ACID MEMBRANE AND PSSAGRAFTED ONTO FEP MEMBRANE AS A FUNCTION OF THE COUNTERION FOR A RELATIVE HUMIDITY OF7%VOH hOH VSSOT n H 2 0 PSS" FEP-PSS PSS" FEP-PSS PSS" FEP-PSS ds FEP-PSSLi + 3458 3453 1631 1643 1041 1041 2.3Na+ 3459 3455 1629 1651 1036 1042 1.4Kf 3460 3454 1631 1655 1035 1044 1.1c s + 3459 3455 1637 1657 1030 1043 0.5" After Zundel, ref.(1)2566 I O N I C HYDRATION I N ION-EXCHANGE MEMBRANESFIG. 7.-Dependence of0 5 10nH 2 0the position of the symmetric stretching vibration bandon the degree of hydration. Symbols as in fig. 4.of the sulphonateatmosphere of 98% humidity. The grafting of these films with PSSA maintains ahigh degree of rigidity as compared with PSS ungrafted membranes and the mem-brane swelling is kept at a relatively low value. Thus, the ratio between the thick-ness of the FEP-PSSA film and the ungrafted FEP film, both films being placed in3500-- -I\ j 3475-5$EFIG. 8.-Dependence of'H 2 0the position of the OH stretching vibration band of thethe degree of hydration.Symbols as in fig. 4.water of hydration oL. Y . LEVY, A . JENARD AND H . D. HURWITZ16402567FIG. 9.- -Dependence of1655 btof hydration on theequilibrium with the laboratory atmosphere, is <1.2. The number of water mol-ecules per equivalent, t ~ ~ ~ ~ ~ , absorbed in the Li’ salt of a FEP-PSSA membranereaches a maximum of 10 at p o , the vapour pressure of pure water at 25°C. Thisvalue is z 1 water molecule below the water content of Li+ salts of PSSA resinscontaining 8% DVB.” An estimate of the lowest amount of water, I Z ~ , ~ , in themembrane follows from our i.r. spectroscopic investigation of thoroughly driedmembranes. On the whole, except for Cs’ and perhaps K+ salts, some watermolecules are very firmly bound to the membrane and cannot be removed.Theirnumber, as it is reported in table 2, reaches an average value of I Z ~ , ~ 2 1 in the caseof Li+ salts.Fig. 5 shows a linear relationship at 25°C between flHZo, the amount of ab-sorbed water molecules, and A,, the integrated do,, absorbances. Such simple be-haviour, which is valid irrespective of the nature of the counterion, can be obtainedTABLE 2.-NUMBER OF WATER MOLECULES ABSORBED IN THE THOROUGHLY DRIED FEP-PSSAMEMBRANEion Lif Naf K+ Csfng,o 0.9 0.5 0.3 0.2568 IONIC HYDRATION I N ION-EXCHANGE MEMBRANES'lH2OFIG. 10.-Dependence of the OH stretching vibration integrated absorbance of the hydration water onthe degree of hydration water in alkaline and alkaline-earth salts of the membrane. Symbols as in fig.4.only after deductions of the peaks due to the polystyrene skeleton in this spectralregion. The hypothesis of constant absorption coefficient of the bending mode,irrespective of ntl,() (thus below n1120 = 3.5 for Ba2+ and 1.7 for Cs') and of thenature of the cation, is justified by the fact that in fig. 5 the linear extrapolation ofthe straight line segment, obtained at equilibrium with the laboratory atmosphere,passes through the origin.As revealed in fig. 4, the conclusion regarding the behaviour of A,,,, the inte-grated vOH absorbances, is quite different. The two slopes of the linear segments canbe interpreted in terms of absorbance coefficients characteristic of the countercatio-nic charge.Furthermore, the plot of A,,, with respect to nHzO in fig. 10 indicates achange in the absorbance coefficients as a function of the water content at lowdegrees of hydration.If the linear dependence of AdOH with nHzO is taken for granted, water absorptionisotherms can be readily derived from combined i.r. and isopiestic measurements.The isotherms recorded in fig. 6 broadly present the classical morphology usuallyencountered in polyelectrolyte systems. A comparison with the absorption isothermin PSSA resins cross-linked with 87; DVB22.23 indicates similar behaviour for K+salts. As already emphasized in table 2, however, one notices with membranes asteep initial rise and a sharp bend. For Li+ salts, H : , , ~ 5 1. Thereafter, a graduallydecreasing sloping curve extends as far as p / p o = 0.40.It follows a moderatelyrising part yielding at p / p o = 1 a smaller water content than in the case ofPSSA-8:l; DVB resins.In order to reach a molecular interpretation of the sorption process of water inthe membrane, it is worth focussing attention on the frequency shift of the principaL. Y . LEVY, A . JENARD A N D H . D . HURWITZ 2569i.r. absorption band maxima with respect to hydration and to the nature of theexchangeable cation.THE SYMMETRIC VIBRATION, Vso3The strength of the ionic interaction between the fixed SO, anion and thecountercation can be evaluated from the position of this absorption band. Con-sideration of the electronic distribution in the direction of one of the s-0 bonds inthe sulphonate ion indicates that the bond energy is larger when the centre ofcationic charge is located on the S-0 bond axis, in contact with the oxygen, thanif the counterion affects all three S-0 bonds together.' The increased probabilityof the existence of the S-0- (cation) structure produces a shift of the maximumof the spectroscopic stretching vibration band to frequencies > 1040 cm-'.* Zundel'has analysed the antisymmetric stretching vibration of SO, in PSSA ungraftedmembranes, but he has not considered the change in the position of the symmetricstretching vibration maximum.On Teflon films, the antisymmetric vibration isobscured by the broad and intense Teflon band at 1200 cm-' ; however, thechanges in the symmetric vibration band become very important.In all alkalinesalts (fig. 7), the shift of the maximum of the vsSoF band with increasing hydrationindicates a weakening of the ion-pair bond strength. At low water content, thestrongest interaction is observed for the Li'-SO, pair, the maximum lying at2 1050 cm-'. At nt120 > 5, however, the binding strength of Li' becomes smallerthan that exerted by other cations. This difference is not too significant because ofthe high intensity of the SO, stretching vibration band.THE STRETCHING VIBRATION B A N D V O HThe vOH stretching vibration band of water at z 3400 cm-' is broad and fairlycomplex. It overlaps with the overtone 280t, of the scissor vibration of H 2 0 at2 3250 cm- ', with the stretching vibration of the CH groups in the benzene ringand partially with symmetric and asymmetric vibrations of the -CH- groups ofpolystyrene.A faint shoulder at z 3600 cm-' can be assigned to the stretchingvibration of the free OH groups pertaining to water molecules linked by theirsecond OH group to an oxygen atom belonging to the sulphonate site or to thewater network contained in the pores of the membrane. The fact that no distinctpeak or shoulder is found at a higher frequency, which could correspond to the OHstretching modes of free water,24 confirms that the number of free water moleculesis negligibly small at any degree of wetness. Due to the weakness of the band at3600 crn-l, the number of free OH from molecules of water of hydration should bevery small, even at the lowest water content.Note that our consideration relies onthe assumption that the absorption coefficient of the free OH group is not too lowcompared to the absorption coefficient of the hydrogen-bonded OH. We have somearguments proving this assertion, since some i.r. spectra of other polyelectrolyticmembrane systems (Nafion) exhibit a sharp band at 2 3600 cm-', associated with asignificant amount of free OH groups in the systems.25The contribution of the OH groups involved in hydrogen bridges is shown by thebroadness and intensity of the band at 3460 cm- An essential cause of the broad-ness of the band is the variation of strength and length of the hydrogen bondswhich link the two OH groups of the water molecules to neighbouring hydrogen907 cni-'.* In hulphonic acid, the S=O and S-0 stretching vibrations are, respectively, located at 1172 an2570 IONIC HYDRATION I N ION-EXCHANGE MEMBRANESbond acceptors.6-2 As has been ~ t r e s s e d , ~ ' - ~ ~ the uncoupling of the vibration ofthe two OH groups produces a broadening of the absorption band. Furthermore,the increase in strength of the hydrogen bond leads to a shift of the maximum ofthe band towards smaller wavenumbers. Hence the behaviour of vOH in fig. 8denotes a continuous decrease in the hydrogen-bond strength with respect to thedegree of hydration. With Li+ salt, one observes a constant value of frequencyabove nIlLo = 5. Furthermore, the association of water molecules with the SO,sites or with each other is progressively enhanced in passing from Cs' to Li'.THE BENDING VIBRATION B A N D 601,1.r.spectral analysis of aqueous solutions of electrolytes has shown that theinteraction between the catiop and the solvent affects essentially the BOH band.33This influence is confirmed in fig. 9 where a drastic difference appears between thebehaviour of the Li' salt and the other alkaline systems. With increasing watercontent, the band maximum with Li' shifts towards larger wavenumbers. WithnHzO > 5, the increment in dOH decreases and the force constant of the H 2 0 scissorvibration reaches at water saturation a value of the same order of magnitude as inice. Contrary to this effect, one observes in dry membranes, following the sequenceNa' < K+ < Cs', a considerable increase in the maximum frequency of hoHabove the value found in ice.This implies, in passing from Na' to Cs', an impor-tant increase in the rigidity of the water hydration molecules. Such phenomenacannot be interpreted as being in a simple relationship with the hydrogen bondingsince the strength of the hydrogen bond decreases along the same ionic sequence.Let us note that such high frequencies have been detected in crystals containingH 3 0 + groups (1700 cm-1)34 or H501 groups (1675 cm-1).35 The addition ofwater to the membrane produces a progressive decrease of the dOH frequenciestowards the value found in liquid water at z 1645 ern-'.MODEL FOR THE HYDRATION OF ALKALINE SALTS OFFEP-PSSA MEMBRANESOn the basis of the spectroscopic results reported here we will attempt to de-scribe a molecular model for hydration of the membrane.The low frequencies ofthe vgH band in the dry membranes (fig. 8) reveal that the sulphonate sites stronglybind the first absorbed water molecules. These molecules possess almost no OHgroups free of hydrogen bonding and thus will bridge two anionic sites, as shown infig. 11. The more water molecules there are absorbed in the membrane, the weakerare the hydrogen bonds formed by the water molecules and the electrostatic inter-action linking cation and anion together. Thus a cationic hydration shell may bemore or less completely built up at the cost of the hydrogen-bond donor proper-ties of the OH groups of the water and at the cost of the strength of the alkalinesulphonate ion-pair association. Conclusions of the same type have been reachedby Zundell for PSSA ungrafted membranes and this author developed the argu-ment that the polarizing effect of the cation on the OH bond is reduced if spreadover several water molecules.Closer examination of our experimental results raises the question as to how thewater molecules are affected by the type of cation fixed in their neighbourhood.Theprofound discrepancy found between the behaviour of the bending vibration doH inthe presence of Li+ on the one hand and the larger alkaline ions on the otherprecludes any simple and general model of cationic hydration. In this respect, thL. Y. LEVY, A. J E N A R D A N D H. D. HURWITZ 257 1IC,Fz In- P S -C - FlCF21nK'PisICF,),- C I - ICF, InFIP SIICFZJn-PS -c -II(CF2)nI\CFJ)n-CdCF2)nFFIG.1 1 . 4 ~ ) Hydration models of the Kf salt of the FEP-PSSA membrane. (b) Hydration model of Lifsalt of the FEP-PSSA membrane.dramatic increase in rigidity of the water molecules following the sequenceNa' < K+ < Cs' and the accompanying decrease in frequency as a function ofhydration is a feature not encountered in PSSA ungrafted membranes. In thesesystems the shift of the scissor vibration under the influence of the alkaline cation isonly 3 cm-'.' Faced with this peculiar behaviour, we will focus our attention firston the large cations and suggest that some clustering of ionic multiplets composedof SO, and alkaline cations should occur, according to the following parameters:(a) their sizes, (b) their mutual repulsions, (c) the vicinity of the anionic sites and (d)the rigidity of the hydrophobic polymeric matrix.A model of stable molecularconfiguration of polyelectrolytes arising through ionic aggregation has already beenproposed in some similar cases. Note, for example, that a dynamic mechanicalstudy of perfluorosulphonic copolymers has yielded different results with Li + salt2572 IONIC HYDRATION IN ION-EXCHANGE MEMBRANEShydrophobic 1 region0 -FIG. 12.-Biphasic molecular configuration in a supermolecular structure.on one hand and the remaining alkaline ions36 on the other. For this system,some experimental arguments based on stress relaxation and small-angle X-rayscattering have led to the suggestion of ion clustering. Such a model of ion cluster-ing, advocated in this context by E i ~ e n b e r g , ~ ~ is not just a local multiplet ionaggregation leading to electrostatic cross-linking in the polymer, but introduceslarge-scale order incorporating a large number of regions of ionic grouping separ-ated by the non-ionic polymeric material.It is difficult to assess how far such amolecular picture of supermolecular ionomeric aggregation can be adapted to ourpresent results. Nevertheless, Hopfinger et al. ’** have conceived a molecular modelfor the case of perfluorinated copolymer membranes of aqueous solvated polymerswith mobile ions and water present in the pores, as depicted in fig. 12. The mem-brane is envisaged as being composed of spherical pores coated by a polymeric skinand drawn together to form supermolecular structures.Inside each pore the anionand countercation interact forming ionic multiplets intermeshed with water. Aninteresting feature predicted by the model is the loss of individual ionic hy-dration shells, the remaining water molecules being accommodated within theinterstices of the ionic network near the sulphonate sites. The rigidity of the poly-meric backbone limits the extension of volume occupied by the ionic multiplets. Itfollows that at low water content the elastic deformation exerted on the packing ofthe hydrogen-bonded water molecules attached to the SO, sites and located in theionic network increases with the size of the countercation.If the relative humidityin the membrane rises, the water molecules, in order to satisfy the free energybalance, will increasingly absorb in the vicinity of the cations and thereby decreasethe lifetime of the cluster [see fig. ll(a)]. This kind of behaviour is in agreementwith the observation of stress relaxation changes in perfluorinated copolymers37 onaddition of water and is also suggested by the position and shift of thevibration recorded in fig. 9.The peculiar behaviour of the small cation Li+ certainly pertains to its ability tointeract strongly with the anion. In some way, the Li+-SO, pair should proveunable to form more than ionic dipole doublets or quadruplets. Thus it will notcontribute to a large-scale cluster organization. The initial water moleculesattached to the SO, sites are therefore strongly polarized by the electric field of the* We thank Dr.L. Bourgeois of Solvay (Belgium) who kindly made this communication known to usL. Y . L E V Y , A . JENARD A N D H. D. HURWITZ 2573nearby cation. This greatly enhances the hydrogen-bond donor properties of theOH group and in addition decreases the energy of absorption of the water molecule[fig. ll(b)]. On increasing the water content in the membrane and as a consequenceof the strong interaction with the cation, a hydration shell is established whichloosens the hydrogen bonds and also enhances the value of the force constant of thebending vibration mode of the water towards that of liquid water, The final settingup of a constant aO1, frequency of the order of magnitude of that of ice, which islarger than the values found for Na' and K', appears consistent with observationson electrolytic aqueous solutions and characteristic of a water structure orderingion with peripheral hydration.33 That the frequencies do not approach a con-stant value at any value of nHzO clearly indicates that even at a relative humidity of80% no definitive hydration configuration of the anion and cation is reached in themembrane.CON CLU S 1 0 N SThe use of i.r. spectrophotometry under isopiestic conditions in conjunction withmore classical techniques, such as microcoulometric titration and Karl Fischertitration, has proved to be a very valuable means of investigating the hydration ofthin ion-exchange membranes.Our results indicate that the presence of a Teflon FEP matrix yields an elasticconstraint exerted on the grafted polyelectrolyte, on the counterion and on waterwithin the pores which might lead, in the case of the large alkali cations, to thesetting up of an ionic cluster configuration. The marked influence of the matrix onthe hydration structures and ionic interactions excludes any attempt, as suggestedby G l ~ e c k a u f ~ ~ and Zundel,' to extrapolate the observation to very concentratedsolutions of electrolytes.The method provides, however, an original insight into themolecular processes responsible of membrane swelling and ionic selectivity.' G. Zundel, Hydration und Intermolecular Interaction (Academic Press, N.Y., 1969).' G .Zundel and A. Murr, Z. Nuturforsch., 1969, 24b, 375.'R. E. Verral, in Water, ed. F. Franks (Plenum Press, 1973), vol. 3, chap. 5.4J. C. Decius, Spectroclzrrn. Acta, 1965, 21, 15.'G. Eisenman, Biophys. J . Suppl., 1962. 2, 259. ' G. N. Ling, A Phy5icul Theorj cf the Living Sttrte (Blaisdell, N.Y., 1962). ' H. P. Gregor, J . Amer. Chem. Soc-., 1951, 73, 642,'S. A. Rice and F. E. Harris, Z. p l t j ~ . Chem. (Frunkfurt), 1956, 8, 207. ' V. S. Soldatov, Russ. J . P h j ~ Cliem.. 1972, 46, 250.'OR. M. Diamond and D. C. Whitney, in Ion Exchange, ed. J. A. Marinsky (M. Dekker, N.Y., 1966),"B. Sansoni. Anyew. Chem., 1963, 75, 164.l 2 J. Ceynowa, Poljmer, 1978, 19, 73.13E D. Schmid and E. Langenbucher. Spectrochim. Actu, 1966, 22, 1621.14S. E. Wiberley, S. C. Bunce and W. H. Bauer, Analjr. Clzem., 1960, 32, 217."R. E. Moynihan. J . Amer. Clzem. Soc., 1959, 81, 1045.I'M. J. Hannon, F. J. Boerio and J. L. Koenig, J . Chem. Phys., 1969, 50, 2829." J. F. Rabolt and B. Fanconi, Mucromolecules, 1978, 11, 740.I'D. H. Whiffen, Spectrochirn. Actu, 1955. 7, 253.I 9 R . R. Randle and D. H. Whiffen, in Molecular Spectroscopy, ed. G. Sell (Inst. Petrol., London, 1955)-'' K . Kiss-Eross, Analytical ItfiareA Spectroscopy, in Comprehenqiue Analytical Chemistry, ed. Wilson'' L. Levy, A. Jenard and H. D. Hurwitz, Anulj't. Chim. Actn, 1977. 88, 377."H. P. Gregor. B. K. Sundheim, K. M. Held and M. H. Waxman, J . Colloid Interjkct) Sci., 1952, 7, 511.13E. Glueckauf and G. P. Kitt, Proc. Roy. Soc. A , 1955, 228, 322.14S. C. Mohr, W. D. Wilk and Ci. M. Barrow, J . Amer. C ' h c m Soc., 1965. 87, 3048.vol. 1.p. 111.and Wilson (Elsevier. Amsterdam, 1976), vol. VI2574 IONIC HYDRATION I N ION-EXCHANGE MEMBRANES2 5 L. Levy, Plz.D. Thesis (Universitk Libre, Brussels, 1979).26G. S. Landsberg and F. S. Baryshanskaya, Izvest. Akad. Nuuk S.S.S.R., Ser. Fiz, 1946, 10, 509.27 Yu.Ya Efimovana and Yu.1. Naberukhin, Mol. Plzys., 1975, 30, 1621.28Yu.Ya Efimovana and Yu.1. Naberukhin, Mol. Plzys., 1975, 30, 1627.29Yu.Ya Efimovana and Yu.1. Naberukhin, Mol. Phys., 1975, 30, 1635.30G. Zundel and E. G. Weidemann, Trans. Faraday Soc., 1970, 66, 1941.32E. G. Weidemann and G. Zundel, Naturforsclz., 1970, 25a, 627.3 3 R. E. Nightingale, in Chenzicul Physics of Ionic Solutions, ed. B. E. Conway and R. G. Barradas (Wiley,34R. Savoie and P. A. Guiguere, J. Chem. Phys., 1964, 41, 2698.35R. D. Gillard and G. Wilkinson, J. Chern. SOC., 1964, 1646.36S. C. Yeo and A. Eisenberg, J. Macromol. Sci., 1977, B13, 441.37 A. Eisenberg and M. King, Ion Containing Polymers and PIzysical Properties and Structure, in Polymer3 8 A. J. Hopfinger, K. A. Moritz and C. J. Hora, communication at the meeting of the Electrochem. SOC.,I. Kampschultre-Scheuring and G. Zundel, J . Phys. Chern., 1970, 74, 2363.N.Y., 19661, p. 87.Physics, ed. R. S. Stein (Academic Press, 1977), vol. 2.Oct. 1977, Atlanta, U.S.A.(PAPER 911812

ISSN:0300-9599    

DOI:10.1039/F19807602558

出版商:RSC    

年代:1980

数据来源: RSC

摘要:

J.C.S. Faraday I, 1980,76,2575-2586Analysis of Electrical Double-layer MeasurementsBY DENVER G. HALL* AND HENRY M. RENDALL?Unilever Research, Port Sunlight Laboratory, Wirral, Merseyside L62 4XNReceived 14th December, 1979We describe methods of presenting experimental electrokinetic and potentiometric titrationmeasurements which are more informative than traditional representations. Simple procedures,based on the application of a theorem recently proved by one of us, make it possible (a) to identifyregions where a simple double-layer model is adequate to fit the results ; (b) to calculate values ofmodel parameters ; (c) to recognise the breakdown of the simple model assumptions ; and ( d ) tosuggest, in many cases, likely causes of the breakdown. A generalisation to more complex systems,with several adsorbed species, is outlined. The application of the theorem to determine the surfacearea of a dispersion is discussed.Electrokinetic measurements and potentiometric titrations are widely used toobtain information about the state of charge at the solid aqueous solution interface.The results are often discussed in terms of the Stern-Grahame model for the doublelayer.Recent publications 1-4 have developed the application of this model toelectrokinetic measurements near the isoelectric point (i.p.). In the simplest case,the following assumptions are made : (1) There is a region (the Stern layer) close tothe surface which is devoid of counterions and coions. (2) The Poisson-Boltzmannequation applies outside the Stern plane.(3) A procedure is available to calculatethe [-potential, at the plane of shear, from electrokinetic measurements. (4) Theplane of shear coincides with the Stern plane.This limiting model will therefore fail if any of the following occur : (a) The planeof shear is displaced from the Stern plane by a distance A, which is not necessarilyconstant. A procedure has been given to identify and in principle to quantify thissituation.2 (b) The basis of the calculations (including the applicability of the PBequation) breaks down. (c) Ions, other than potential determining ions, are (speci-fically) adsorbed within the Stern plane.However,to calculate, from titration results, the charge density at the interface, it is necessaryto know the surface area, A, of the sample and the point of zero charge (P.z.c.),which is less easily obtained than is the i.p.The choice of experimental techniquedepends on the nature of the sample and on the information required.Although measurements in the immediate region of the i.p. are well under~tood,~'~the analysis of results becomes tedious and unsatisfactory at higher charge densitieswhere the simple model assumptions are increasingly likely to break down and wherein particular the specific adsorption of counterions may be expected. A simplecriterion to establish limits, outside which the assumptions are clearly not applicable,was proposed recently.' The purpose of the present paper is to explore the sensitivityof the proposed test and to establish whether, from the form of the deviations, areasonable diagnostic inference about the likely cause of the breakdown may beobtained.For potentiometric titrations, only assumptions (1) and (2) apply.t Present address : Paisley College of Technology, High Street, Paisley, Renfrewshire PA1 2BE.2572576 ANALYSIS OF DOUBLE-LAYER MEASUREMENTSTHEORYConsider a simple double-layer model consistent with assumptions (1)-(4), in thepresence of a z : z supporting electrolyte and with a surface excess rl of potentialdetermining ion.In terms of rationalised units we haveandgo = zleorl (1)2 k T a sinh (-). zleO$d2kT - g d = -z1eoWe definePl = In'l-z,eo$a (3)(4)wherepy = pF( T ) + kT In n'; .The theorem proved in ref. (5) shows that p 1 -@ is the same for all points with thesame rl.Thus, at constant C T ~ , ~where X is a potential determining ion and the constant is a function of do. Underthese conditions, therefore, $d follows a " Nernst " relation with pX. An equivalentstatement is that, over the entire experimental range of pX and ionic strength ( I )where assumptions (1)-(4) hold, plots of go (or ad) against ,ul as defined by eqn (3)should be congruent.When the surface potential tjo follows a Nernst relation with pX, eqn (5) has theformZleOtjd = const -2.303kT pX ( 5 )2.303 kTPX2.303kTwhere pX" is the value of pX at the isoelectric point and K the integral capacity of theStern layer. For a surface with two ionising groups of the type described in ref.(3),we havePX (7)$d=[- 2.303kTz1eowhere Y = [( 1 - O-)/O-][O+/( 1 - O,)] is constant for constant do.It follows that experimental results which have already been shown to be consistentwith the Stern model will inevitably conform to eqn (5). We may therefore use theseresults to demonstrate the possibility of applying eqn (5) for a preliminary assessmentof experimental results and to examine the relative merits of eqn (5) and of the con-gruence test.RESULTS AND DISCUSSIONCOMPARISON WITH EXPERIMENTFig. 1 shows (interpolated) values of 5 against ApCa for CaCO, and Ca,(PO,),at constant values of dd, with theoretical lines as given by eqn (6). Fig. 2 shows asimilar plot for a nylon The proposed test would have been particularly usefulin this case, where it is not immediately obvious that the results should or could beanalysed on the basis of a simple double layer model.Fig. 3 shows the treatmen3020100>E2ii -10-20-304aD . G . HALL AND H. M. RENDALL 2577I 11-2 -1 0 1 2ApCaFIG. l.-Plots of {-potential against pCa, at constant Ud for CaCO,(O) and Ca3(P04)2 (O)." Solidlines given by eqn (6) with K = 40 pF cm-2 and - Ud = 0.5 (a), -0.5 (6) and - 1.0 pC cm-2 (c).PHFIG. 2.-Plots of (-potential against pH, at constant (Td, for nylon 5 0 1 . ~ Solid lines given by eqn (7) withK = 25 pF cm-2 and - Ud = 0.4 (a), 0.3 (b), 0.2 (c) and 0.1 pC cm-2 ( d ) . Appropriate values of Ywere obtained from the data in ref. (3).1-82578 ANALYSIS OF DOUBLE-LAYER MEASUREMENTSapplied to electrokinetic results with AgI (sols were prepared by precipitation fromKI/AgF, dialysed not less than 24 h and used within 2-3 days.Particle sizes weretypically 50 nm. The supporting electrolyte was KN03).7 The best fit overall isobtained with K = 25 pF cm-2, although a slightly better fit at low charge densitiesis obtained with K = 27 pF cm-2, as was calculated by Peterson from the ionicstrength dependence of the quantity Ns-l.l* In principle, therefore, the presentmethod of representing the results may given information on the value of K =f(o,)while assumptions (1)-(4) are valid. Deviations from the proposed plot set in athigh la1 and at high ionic strengths (all experimental points at ionic strength 0.15 ~ m - ~and for ladl > 1 pC cm-2 at 0.1 mol dm-3).These conditions represent a severetest of the theoretical assumptions and, indeed, of the experimental technique. Thedeviations are, however, in a direction qualitatively consistent with the effect ofspecific adsorption of counterions. Fig. 1-3 demonstrate the application of theproposed test to ascertain whether a simple Stern-layer model can be expected tofit the results (see especially fig. 2) and to identify regions where the simplest modelwill clearly not apply (fig. 3). To confirm this usefulness of the proposed plot,however, it is desirable to show more clearly that experimental results to which thesimple double layer model is not applicable will unequivocally fail the test. Electro-kinetic measurements for quartz or silica interfaces are particularly appropriate forthis purpose.It has already been shown that results for silica do not conform90807060SO > s u4030201c2 3 4 5PAgFIG. 3.-Plots of &potential against pAg, at constant ad, for AgI sol.’ Solid lines given by eqn (6) withK = 25 pF Cm-’ and - ad = 3.0 (a), 2.5 (b), 2.0 (c), 1.5 (d), 1.0 (e) and 0.5 pC cm-’ (f). Ionicstrength (0), lo-’ (@), 5 x lo-* (a), 10-1 (a) and 1.5 x 10-1 mol dm-3 (0)D . G. HALL AND H . M . RENDALL 2579to a stability test based on assumptions (1)-(4). The [-potentials for the samesample *, are shown in fig. 4 as a function of pH. Although reasonably linear plotsof [ against pH at constant (Td are obtained for a number of values of od, the slopesare = 33 mV per pH unit, clearly far removed from the " Nernst " value implied byeqn (5).The proposed test therefore correctly identifies the breakdown of the modelrepresented by assumptions (1)-(4) for a system where this failure has been established.The alternative " congruence " test is shown for these samples in fig. 5-7, whereb d is plotted as a function of pX* = pX+zle()$d/2.303kT [eqn (3)]. This procedurehas the advantage of using directly all the experimental results. Fig. S demonstrates,perhaps more convincingly than fig. 2, that the ionic strength-dependence of themeasured nylon mobilities may be attributed to diffuse layer effects. On the otherhand, fig. 3 demonstrates more graphically than fig. 6 the deviations from the simplemodel of the results for AgI at high ladl and ionic strength.The silica results (fig. 7)-100PHFIG. 4.-Plots of S-potential against pH at constant Ud for silica. (a), Ud = 0.2 pC cm-2, slope =31.7mV ; (b), ad = 0.4 p C cm-2, slope = 32.7 mV ; (c), Ud = 0.5 pC cm-2, slope = 31.8 mV ; (a), Ud =1.0 pC cm-2, slope = 34.8 mV2580 ANALYSIS OF DOUBLE-LAYER MEASUREMENTSfail the congruence test under all experimental conditions, confirming the inapplica-bility of the simple model used here. The same general features are observed in theelectrokinetic properties of quartz.” Note that for silica and quartz the variationof od with ionic strength is in the opposite direction from that which would be expectedfor simple specific adsorption of counterions, or from a constant displacement A ofthe plane of shear.A detailed discussion of the nature of the double layer on silicais outside the scope of the present paper. However, considering assumptions (1)-(4),the form of the curves in fig. 7 appears to imply either the presence of coions withinthe plane of shear, or a breakdown of the model for calculating 5 from electrokineticmeasurements. In the absence of titration curves for the same samples as wereused for the electrokinetic measurements, the Naf adsorption results of Li andde Bruyn lo may give some indication of the plausibility of these suggestions. Onthe assumption that the Na+ adsorption was governed purely by the Poisson-Boltzmann equation, these authors lo deduced double-layer potentials for quartz atmol dm-3 agreeing (except at extremely high potentials) with themeasured electrokinetic potentials.This, therefore, provides support for the validityof the [-potential calculation and eliminates one possible source of the “non-congruence” of the results (fig. 7) at low ionic strength. These Na+ adsorptionresults, unlike (hydrogen) potentiometric titrations, do not allow us to eliminate thepossibility that coions, but not counterions, are found within the shear plane. Athigher ionic strength ( 2 mol dm-3), the potentials deduced from Na+ adsorptionwere of substantially higher magnitude than those from electrokinetics. Here,andI I I II I I I4 6 8 10 -1 .c2pH + zleO 512.303 kTFIG, 5.-Plots of Od against pH* for nylon Ionic strength lo-* (o), (O), 5 x (A) andmol dm-3 (0).Solid line given by eqn (8) with KA = 8.6 pF cm-2therefore, at constantkinetic measurementsor specific adsorptionconsistent explanationsamples.D. G . HALL A N D H . M. RENDALL 258 1ionic strength, intercomparison of " titration " and electro-would suggest outward displacement of the plane of shear,1°of counterions. Hence there is at present no obvious self-. of the failure of the congruence condition for silica/quartzDEDUCTION OF DOUBLE-LAYER MODEL PARAMETERSRearranging eqn (6) we haveHence the value of K may be obtained directly. If we represent the slope[da,/d(pX*)] by the symbol S, the relation with the slope d@,/d(pX), at the i.p.,t t 13 G 5pAg+ (z1e0</2.303 kT)FIG. 6.-Plots of Od against pAg* for AgI.' Ionic strength low3 (O), (@), 5 X lod2 (Q), lo-' (a)and 1.5 x 10-1 mol dm-3 (0).Solid line, eqn (8) with K = 25 pF cm-22582 ANALYSIS OF DOUBLE-LAYER MEASUREMENTSpreviously defined as sl, is given byKEN _ - - 1-- NS1 S (9)where N, the “ Nernst ’’ factor, is defined as -2.303kT/z,eo.However, unlike sl, which is a limiting slope where $d 4 0 and which varieswith ionic strength, the valfie of S is a constant for all points where assumptions(1)-(4) hold, for a “ Nernst ” surface with a constant double-layer capacity K. Thismakes the evaluation of the slope and hence of K simpler and more reliable. IfK = f(ao), or if a low density of ionising sites gives rise to deviations from “ Nernst ”beha~iour,l-~ the prescribed plots will be non-linear, but congruent.Hence theseeffects may be distinguished from specific adsorption. Clearly, also, the sameapparent capacity KA for a “non-Nernstian” interface is obtained at the i.p. asfrom the Ns-1 method.In confirmation of this, a linear regression analysis gives KA = 8.52 pF cm-2(r = -0.963) for nylon (fig. 5, 17 points). This compares well with the reportedvalue of KA = 8 6k0.2 pF cm-2. For AgI (fig. 8, including electrokinetic results2.0 -1.5 -NI Eg 1.0- \ Y80 s -14 6 8pH+ (z1e0C/2.303 kT)FIG. 7.-Plots of ad against pH* for ~ i l i c a . ~ Ionicstrength (a), (0) and mol dm-3 (A)D. G . HALL AND H . M. RENDALL 2583from fig. 6 and titration resu1ts)ll K = 26.4 pF cm-2 (Y = -0.977), below the i.p.The titrated charge densities at pAg values above the i.p.deviate from the linerepresented by K = 26.4 pF cm-2, but apparently remain congruent. This effectwould therefore be attributable to changes in the value of the double layer capacityK rather than to specific adsorption of supporting electrolyte ion.21NI0 3 . Yb"I-1- 2I II I- 1 0 1pAg+ (z1e05/2.303 kT)2FIG. 8.-Plots of adagainst pAg* for AgI. Resultsfromtitration '' (open symbo1s)andelectrokinetics '(filled symbols) at ionic strength (0) and lo-' mol dm-3 (0). Solidline, eqn (8) with K = 26.4 pF cm-2. Bars indicate spread of reported titration results at constantpAg or constant cro.Finally, we consider the case where the shear plane is displaced a constant distanceA into an otherwise undisturbed diffuse layer.Plots of the electrokinetic chargeagainst the pX* evaluated at the plane of shear will no longer be congruent. How-ever, near the i.p. the slopes of these plots will follow a relation of the same form aseqn (9) with the limiting slope s[dc/d(pX)], defined previous1y.l- The slope s isgiven by(0), lo-* (A), 5 xs = s1 exp (-.A). (10)In this situation the present method gives no additional insight into the interpreta-However, because the plots are at least partially linearised, tion of the measurements2584 ANALYSIS OF DOUBLE-LAYER MEASUREMENTSevaluation of the initial slopes and hence of A may be more reliable than from adirect [-potential plot.Some reasonable deductions from the form of the '' congruence plot " for electro-kinetic measurements are summarised in table 1.A quantitative analysis of theresults would normally be required to distinguish the alternative explanations of agiven observation. If, however, potentiometric titration results are available for thesame sample, significant qualitative differences may be observed. For example,outward displacement of the plane of shear would not influence the results of potentio-metric titrations, nor indeed would the breakdown of other assumptions, required inthe calculation of ( from electrokinetic measurements, about the response of thesystem to a non-equilibrium situation. On the other hand, specific adsorption ofcounterions would affect the " congruence plot " for titration results, but in theopposite sense from that quoted for electrokinetics. The final case in table 1,observed here with silica but probably characteristic of a range of materials, is notfully understood and clearly merits further consideration.TABLE 1 .-CONGRUENCE PLOT FOR ELECTROKINETIC MEASUREMENTSobservation inference(1) linear, congruent plot results will be fitted adequately with a" Nernstian " surface potential and a constantcapacity K(2) non-linear, congruent plot (4 K = f(odor (b) non-Nernstian Ic/o : pX relation(a) specific adsorption of counterions(b) (outward) displacement of plane of shear or(a) specific acisorplion of coions(b) breakdown of assumptions for calculating j oror (c) ?SURFACE AREA DETERMINATIONFor any surface where assumptions (1)-(2) hold [or (1)-(4) in the case of electro-kinetic measurements] the approach outlined above suggests a number of possibleprocedures for calculating the surface area of the sample.( A ) If both electrokinetic measurements (giving a surface charge density) andpotentiometric titrations (giving a total surface excess charge) are available, calculationof a surface area is obviously trivial.This surface area determination should bereliable provided that it is confined to a region where the electrokinetic measurementssatisfy the congruence condition.( B ) If only potentiometric titration results are available, but the sample has aclearly defined and identifiable P.z.c., the area may be determined as follows. Thecongruence condition requires that, for a given surface excess charge N', per unitweight of sample, the slope [d(ApX'k)/aN",l, is independent of ionic strength.ProvideD . G . HALL AND H . M . RENDALL 2585that the slope is evaluated through the P.z.c., the Poisson-Boltzmann equation maybe linearised to provide the simple relation:lim( aApX -) = const + Z;f?;N:+O aN; K E 2,303 k TKEATo obtain the best value for the limiting initial slope of the titration curve, it isconvenient to plot ApX/N; against ApX or N,". Within the linear region of thePoisson-Boltzmann equation, ApX/N," is a constant for a given ionic strength,provided that K and d$,/d(pX) are constant. The surface area of the sample maybe calculated from the gradient of the linear relation, eqn (1 l), between the initialslopes and I/KE.I t is easy to show that the ratio of the gradient to the intercept ofeqn (1 1) gives the apparent capacity KA of the Stern layer which would be obtainedfrom the Ns-l procedure described previously lm2 (to which this present plot isdirectly analogous).(C) For a surface where the P.Z.C. is ill-defined or not experimentally accessible(e. g., a surface containing only carboxyl- or sulphate-type charged groups), procedure(B) may not be applied. Here we make use of the fact that, at constant uo (i.e.,constant NE), the quantity pX+~~e~$~/2.303kT is a constant. In the linear regionof the Poisson-Boltzmann equation, therefore,pX+ [zfe:N;/2.303kT~&A] = constant. (12)When measurements of N," may be obtained at sufficiently low charge densities, thearea may be calculated directly from a plot of pX against 1 /rc at constant N,".Other-wise an initial value of A may be calculated at the lowest accessible N,". This estimatemay then be refined by converting the titrated charges to charge densities and adjustingthe value of A to produce the best congruence plot. An error of at most 10 % in thesurface area is sufficient to generate a clear systematic deviation from congruence.In the event of significant specific adsorption of counterions, the curves could not bemade congruent by any choice of area A . In this case, the appropriate value of Awould be that giving congruence at low charge densities and ionic strengths. Oncethe area is known, the proposed tests may be applied to identify the onset of specificadsorption of counterions.With titration curves having a common intersectionpoint, indicating no significant specific adsorption near the P.z.c., the tests may bebased on the area obtained directly from procedure ( B ) .For a totally uncharged surface, the surface area could be obtained by method (C),provided that species 1 (the potential determining ion X) is replaced by a singlestrongly adsorbing ion.EXTENSION TO MORE COMPLEX SYSTEMSA similar treatment to the above can be applied when there are several ionic speciespresent inside the Stern plane (SP). To show this we recall that, according to thetheorem on which the treatment is based, the quantity dL given bydL = XTi dpiiis an exact differential where ( I ) the summation is over all independent ionic speciesfound on both sides of the SP; (2) the p i are defined byp i = @(T, p ) + kT In i$' - zieoi,hd 2586 ANALYSIS OF DOUBLE-LAYER MEASUREMENTS(3) the Ti are the amounts adsorbed in the inner regions of the double layer andsatisfy the expressionEzieorl+ad = 0.(15)iLet species 1 be one of the potential determining or specifically adsorbed species andletWe may use eqn (14) to substitute for the dpu, in eqn (1 1) which then becomesIt is apparent from this equation that at constant T and ei, pl is a function of bdonly and vice versa. Thus when ions of the supporting electrolyte are absent fromthe inner regions of the double layer it follows that at constant T and Oi graphs of bdagainst [kT In n1 - zle,t,hd] for different ionic strengths should be superimposableand that at constant T, 81 and b d graphs of kT In nl against z,e,t,hd obtained from dataat different ionic strengths should be linear with unit slope.It is clear from the above that both the “ Nernst plot ” and the congruence plotcan be applied to systems in which several ionic species are believed to occur on bothsides of the SP in exactly the same way as when there is only one such species, providedthat the for all species concerned are held constant.This condition correspondsto the quantities (n:)>z’/(n;)zi being constant for all ij pairs.CONCLUSIONSTwo simple procedures are outlined for analysing electrokinetic or potentiometrictitration measurements. The suggested plots offer advantages over other methodsfor deriving model parameters from experimental results. Furthermore, a directmeasure is given of the range of validity of the model assumptions, together with auseful diagnostic indication of the most likely source of the breakdown. There isno problem, in principle, in generalising the treatment to take account of specificallyadsorbing species.We thank Dr. G. C . Peterson for supplying us with an extensive set of unpublishedmeasurements with AgI.A. L. Smith, in Dispersions ofpowders in Liquids, ed. G. D. Parfitt (Applied Science, London,2nd edn, 1973), p. 93.A. L. Smith, J. Colloid Interface Sci., 1976, 55, 525.H. M. Rendall and A. L. Smith, J.C.S. Faraday I, 1978,74, 1179,H. M. Rendall, A. L. Smith and L. A. Williams, J.C.S. Faraday I, 1979, 75, 669.D. G. Hall, J.C.S. Faraday 11, 1978, 74, 1757.T. Foxall, G. C. Peterson, H. M. Rendall and A. L. Smith, J.C.S. Faraday I, 1979, 75, 1034.T. Foxall and G. C. Peterson, unpublished results.D. G. Hall and M. J. Sculley, J.C.S. Faraday II, 1977, 73, 869.G. R. Wiese, R. 0. James and T. W. Healy, Disc. Faraday SOC., 1971, 52, 302.J. Lyklema, results compiled from (a) E. L. Mackor, Rec. Trav. chim., 1951, 70, 763; (b)J. A. W. van Laar, Thesis (State University of Utrecht, 1952) ; (c) J. Lyklema, Trans. FuradaySOC., 1963, 59, 418 ; ( d ) B. H. Bijsterbosch, Thesis (State University of Utrecht, 1965) ; (e)B. H. Bijsterbosch and J. Lyklema, J. Colloid Sci., 1965, 20, 665.lo H. C. Li and P. L. de Bruyn, Surface Sci., 1966,5,203.(PAPER 9/1984

ISSN:0300-9599    

DOI:10.1039/F19807602575

出版商:RSC    

年代:1980

数据来源: RSC

摘要:

J.C.S. Faraday I, 1980,76,2587-2603Some Reactions at a Mercury@) Sulphide PhotoanodeBY R. STEPHEN DAVIDSON"? AND CHARLES J. WILLSHERDepartment of Chemistry, University of Leicester, Leicester LE 1 7RHReceived 9th July, 1979The pho t oelec t rochem ical react ions of pigmentary mercury(I1) s ul p hide have been investigatedusing electrochemical cells having a platinum electrode coated with the sulphide. E.m.f. and currentmeasurements show that the sulphide behaves as an n-type semiconductor and when sodium nitrateis used as electrolyte the photoanode is stable. Use of other electrolytes can lead to solubilisationand a change of colour of the sulphide. The darkened form of the sulphide so produced is morephotoreactive than the red form. The sulphide photoassists the electrolysis of water and this isrationalised on the basis of an energy level diagram drawn up from experimental data.We have previously described the preparation of platinum electrodes coveredwith pigmentary titanium dioxide and the photoelectrochemical properties of suchelectrodes.The ease of preparation of such electrodes and the fact that a pigmentarysample of the semiconductor rather than the more expensive single crystal can beused to carry out photoelectrochemistry suggested that techniques developed in thiswork could be applied to other semiconductors with the hope that a material maybe found which is suitable for transducjng visible light into electrical energy or asource of fuel. The requirements for such a material are (a) it absorbs visibleradiation, (b) it is photostable and (c) it either has a good power output or electrolysesor photoassists the electrolysis of water.We now report upon the photoelectro-chemistry of mercury(I1) sulphide which fulfils, in part, the above criteria.EXPERIMENTALMercury(@ sulphide (vermilion) (Koch Light) was used as received. Other sources ofmercury(i1) sulphide included B.D.H., Fisons and May and Baker.In nearly every case the sulphides showed a negative photo-e.m.f. and displayed a photo-anodic current. Meta-cinnabar (Alfa Products) was used as received.Electrolyte solutions were made up using de-ionised water and the salts were of the highestpurity available.The loading of mercury(I1) sulphide is approximately 0.025 g. Experiments were carriedout with a single compartment cell fitted with a cover and provision for ckoxygenation of theelectrolyte by purging with nitrogen.The cell contained three electrodes, the coated platinumelectrode, a platinum counter electrode and a reference electrode (saturated calomeleIectrod.e, s.c.e.). Potentials were measured with a Philips high impelmcc voltmeter(niodel PM 2434) and currents with either a Heath polarography module (EUA-19-2) or aWenking potentiostat (LB 75L).The lightbeam was focused on one face of the covered platinum gauze. For studying the spectralresponse of the electrodes, the light was filtered by use of broad band interference filters(Balzers) and for studying the effect of light intensity, neutral density filters were interposedbetween the cell and light source.t Present address : Chemistry Department, City University, St.John Street., London EC4 4PB.2587The magnitude of these values varied from source to source.The prcparation of the electrodes has been previously described.2*The light source and solution filter system have been previously de~cribed.~In these cases, water formed the heat filter2588 REACTIONS AT MERCURY(I1) SULPHIDE PHOTOANODEIrradiation of suspensions of mercury(I1) sulphide in electrolytes was carried out with aPyrex jacketed water-cooled 125 W medium pressure mercury lamp (Hanovia).Analysis of electrolyte solutions for solubilised mercury was carried out with a Perkin-Elmer 360 atomic absorption spectrometer. The lowest reliable detection limit was7.5 p g ~ r n - ~ .RESULTSRED MERCURY(II) S U L P H ~ D E I N SODIUM NITRATE SOLUTIONIrradiation of platinum electrodes coated with red mercury(I1) sulphide produceda negative photo-e.m.f., the magnitude of which depended upon the source of thesulphide.Usually it was between - 100 and -200 mV. The photoresponse wassimilar to that reported for titanium dioxide, i.e., on commencement of irradiationthe e.m.f. rapidly built up to its maximum value which was maintained until termina-tion of irradiation when it slowly decayed until the dark e.m.f. was attained. Themagnitude of the photocurrents and the onset of relative anodic photocurrentsdepends upon the source of the sulphide. Fig. 1 illustrates photocurrent againstapplied potential at pH 7.The onset of photocurrent usually lay between -0.1 5and -0.35 V (against s.c.e., pH 7) ; -0.2 V was taken to be the flat-band potential.The effect of change in pH of the electrolyte upon the potential required to causecurrent to flow and upon the photo-e.m.f.s is shown fig. 2 and 3. The onset ofrelative anodic photocurrents, the dark potential and potential upon illurnination(against s.c.e.) all vary by z -0.06 V per pH unit although the relative photo-potential is pH-independent between pH 2 and pH 12. In very acidic and veryalkaline electrolytes, dark currents are large and relative photopotentials diminished.The spectral response of the photovoltage and photocurrent are shown in fig. 4.The effect of light intensity upon photovoltage and photocurrent was determined byinterposing neutral density filters between the light source and the cell and the resultsare shown in fig.5.In order to test for photoinduced decomposition of the sulphide, electrodes wereirradiated for several hours under bias and the electrolyte analysed for mercury ions(Hg2+) by means of atomic absorption spectroscopy. Within the pH range 2-12no solubilisation could be detected. electrons were passed anodicallythrough a sulphide electrode on illumination. If two electrons can solubilise onemercury atom, according to1.2 x 10'HgS 4 Hg2+ + S +2ethis would lead to M 2 x g of mercury in solution. The volume of electrolyteemployed was 150 cm3, so theoretically a mercury concentration of 13 pg ~ m - ~should be detected.A value of 0-1 , u g ~ m - ~ was observed. Usually the mercurysulphide electrode showed a 1-2 % weight loss after irradiation. This small loss isattributed to handling.Occasionally, irradiation of red mercury sulphide electrodes in nitrate solutioncaused the sulphide to change to a brown-black col~our. This colour change isdiscussed in more detail later on. When an anodic bias was applied to the sulphideelectrode gas bubbles were sometimes formed at the sulphide and platinum counterelectrode. Gas evolution was more noticeable if the sulphide electrode had becomedarkened on illumination. Gas evolution was also favoured by increasing the pHof the solutions. Since varying the sodium nitrate concentration between 10 tomol dm-3 had little effect upon the photocurrents, the photocurrent is attributedto the photoassisted oxidation of hydroxyl ions at the surface of the sulphide anR .S . DAVIDSON A N D C . J . WILLSHER 2589Ix x-XXXX XI I I I I I I I- 0.8 -0.4 8.0 0.4applied po tential/V against s.c.e.FIG. 1.-Relative photocurrent against applied potential for red HgS in neutral 0.1 mol dm-3NaN03 (N,-purged, unbuffered). The bias setting is manually adjusted and the dark current allowedto settle before illumination. The flat-band potential is considered to be -0.15 to -0.35 V, fromthe photocurrent onset.0 4 8 12PHFIG. 2.-Approximate onset of anodic relative photocurrent at red HgS against pH of N2-purged0.1 mol dm-3 NaN03 electrolyte.The pH is varied by NaOH or HN03 addition and the voltageat which no net photocurrent flows was determined by sweeping the potential at 1 mVs-l. Thegraph has a slope of M -0.07 V pH-'. The form of the photocurrent-bias plot at all pHs investi-gated is similar to that shown in fig. 3. Reference slope at -0.059 V pH-l2590 REACTIONS AT MERCURY (11) SULPHIDE PHOTOANODEPHFIG. 3.-Variation with electrolyte (0.1 mol dm-3 NaNO,) pH in dark potential (0) and potentialupon illumination ( X ). The relative photovoltage is pH-independent although its magnitude variesfrom sample to sample. Reference slope at -0.059 V pH-l.- 6 0 1 Ic- -60 -I-20 -400 500 6 00wavelength/nm400 I 500FIG. 4.-Wavelength response of photocurrent (solid line) and photovoltage (dashed line) for redHgS in 0.1 mol dm-3 NaN03 obtained by irradiation with 1.8 kW xenon lamp with various broadband interference filters and 11 cm of H20 in the light pathR .S . DAVIDSON AND C . J . WILLSHER 259 1the reduction of protons at the counter electrode. The minimum applied potentialto achieve gas evolution at the platinum counter electrode (with respect to thesaturated calomel electrode) was found to be +0.2 V.RED MERCURY@) SULPHIDE I N ELECTROLYTES OTHER THANSODIUM NITRATEIt was found that darkening of red mercury(I1) sulphide leads to an increase inphotocurrents. The use of electrolytes other than sodium nitrate was investigatedto see if the darkening process was affected by ions in the electrolyte. Use of thefollowing salts as electrolytes induced darkening : potassium ferrocyanide, iodate,iodide, thiocyanate, bromide, chloride, fluoride, cyanide and lead nitrate.Potassiumsulphate, like sodium nitrate, behaved capriciously. The use of many electrolytescaused solubilisation of mercury, e.g., potassium oxalate, ferrocyanide, bromide,iodide, iodate, thiocyanate, sodium thiosulphate, hydrogen-phosphate, stannouschloride, ferrous sulphate, manganous nitrate, thallous nitrate and cerous nitrate.Solubilisation occurred when mercury sulphide electrodes were irradiated in theelectrolyte solution and also when suspensions of mercury sulphide in these electro-lytes were irradiated (see table 1). When some ions having Eo values betweenz -0.5 V and + 1.0 V (against s.h.e.) are present as electrolyte ions, they competewith hydroxyl ions and/or water for reaction with the photogenerated holes.Largerphotocurrents were noted in these electrolytes than for a nitrate electrolyte of thesame pH. This is taken as evidence of successful Competition for holes. Someresults are presented in table 1.Usually the photo-response of the mercury sulphide electrode was independent of the type of electrolyteand was similar to that for 0.1 mol dm-3 sodium nitrate provided the pH of theIn all cases the electrolyte solutions were 0.1 mol dm-3.(a) dintensity (arb. units) (b) light intensity (arb. units)FIG. 5.-(a) The dependence of relative photovoltage of red HgS in 0.1 mol dm-3 NaN03 on lightintensity. At high intensity the photopotential saturates.(b) The dependence of relative photo-current (at 0.0 V against s.c.e.) of red HgS in 0.1 mol dm-3 NaN03 on light intensity2592 REACTIONS AT MERCURY(II) SULPHIDE PHOTOANODEsolution is taken into account. Potassium ferrocyanide and thiocyanate proved tobe the exceptions in that the use of these compounds as electrolytes led to the pro-duction of positive photo potentials (see table 2). That this peculiar behaviour isdue to the electrolyte and not to a change in the mercury sulphide was shown by thefact that when the electrodes were removed from the ferrocyanide or thiocyanatesolutions, washed with distilled water, dried and then irradiated in 0.1 rnol d ~ n - ~sodium nitrate solution they showed the normal photoproperties.(a) [KI]/niol dm-350 0400mI300DuLon 1 c200100000000 0000I0 ' 0XXXXXx xXXI x I 1 I I I I I I n20 40 60 80FIG.6.-(a) Solubilisation of mercury upon irradiation (solid line) and in the dark (dashed line) fromred HgS in various KI concentrations. Irradiation is by a 125 W medium pressure Hg lamp(Hanovia) and the suspension is 2 g of HgS in 500 cm3 of stirred KI (open to the air), for 5 h.(b) 2 g of red HgS in 500 cm3 of 0.1 mol dm-3 KI ; solubilised mercury from irradiation of stirredsuspension with 125 W Hg lamp against time. The solubilisation in the dark was FZ 6 pg ~ m - ~ foreach run. The solubilisation rate is M 4 p g C M - ~ h-ITABLE SOME RESULTS FOR RED HgS ON PLATINUM IN SELECTEDThe effective irradiated area is 2.25 cm2 and the light intensity m 2.5 xrelativephoto-relative currentphoto- at 0.0 V photocurrent no.of electronsb [Hg]electrolyte (aqueous voltage against onset passed anod.ically electrolyte0.1 mol dm-3) pHa /mV s.c.e./A against s.c.e. upon irradiation / p7,77110.711.3599.512.27- 75- 200- 200- 125- + 80+ 120- 180- 150- 170-115- + 5- 200+ 70 + 15+3+4- + 100- 1+2f 5+4+6 + 50+2- 0.35- 0.25- 0.25(negative+ 0.05+ 0.05of - 0.45)- 0.65-0.13- 0.35- 0.25- 0.75- 0.252 . 6 ~5 . 8 ~ 1 0 I 81.2x 10185.5 x 101 74.5 x 10' anoc'.ically5.8 x 10'' cathodically2.1 x anod.ica1ly2.6 x loi6 cathodically2 . 6 ~ 1 0 1 73.1 x 10178 .5 ~ 1 O I 68.9 x 101 73 . 2 ~ 10"w 1 0 I 8mmmG3G3a Electrolyte is unbuffered. This column does not include electrons contributing to any currentHgS in 500 cm3 of electroIyte, stirred and irradiated for 240 h by a 125 W medium pressure Hg lamp.corresponding suspension stirred in the dark for 240 h. Readings below reliable detection h i t oTABLE 2.-cURRENTS AND POTENTIALS FOR A RED HgS ELECTRODE IN AN ELECTROLYTE OF VARIEDindicates intial dark potential and " after " means dark potential afterpotassium potential upon relative relative onset potentialt hiocyanate illumination photo- photocurrent of photo-/mol against s.c.e. against s.c.e. /mV against s.c.e. against s.c.e.- 270 (black)1 .o +255 (red) + 290 + 460 - 9.5 > +0.3 V- 170 (black)The sign of the photoeffect depends upon [SCN-] and mercury solubilisation is noted for [KSCN]coccentration dark potential /mv potential /FA at 0.0 V current/V-5 + 155 (red) - 150 + 120 - 3.5 w +0.2s 0-1 + 255 (red) - 105 + 35 - 0.8 + 0.05+235 (red)- 60 (darkened) - 125 - 65 + 2.4 -0.12s 0-3 + 145 (before) - 140 - 150 + 3.2 w -0.1 + 10 (after)10-4 + 180 (before) - 105 - 175 + 1.3 w -0.1 + 70 (afterR .S . DAVIDSON A N D C. J . WILLSHER 2595TABLE 3.-vARIATION IN SOME PARAMETERS (IRRADIATION TIME, [KI]) FOR DARKENING REDHgS AND SOME RESULTS OF DARKENED ELECTRODESOptimum blackening occurs for [KI] = 0.1 mol dm-3 and time of irradiation = 30-180 min,with the HgS held at 0.0 V against s.c.e.~~ _ _ _ ~ _ _ _ ~ ~ _ _ _ _relative relativephotocurrent photocurrentsrelative blackened of blackenedphotopotential form/pA, form/pAconditions for blackening of blackened at 0.0 V bias at +0.4 V biasred HgS form/mV against s.c.e.against s.c.e. remarksirradiated" for 30 rnin atopen-circuit in 0.1 mold~n--~ KIirradiated" for 30 min at0.0 V (against s.c.e.) in0.1 mol dm-3 KIirradiated" for 60 rnin at0.0 V (against s.c.e.) in0.1 mol dm-3 KIirradiated" for 180 min at0.0 V (against s.c.e.) in0.1 rnol dm-3 KIirradiated" for 300 min at0.0 V (against s.c.e.) in0.1 rnol dm-3 KIirradiated" for 30 min at0.0 V (against s.c.e.) inmol dnr3 KIirradiated" for 30 rnin at0.0 V (against s.c.e.) inlo-' rnol dm3 KIirradiated" for 30 rnin at0.0 V (against s.c.e.) in1.0 mol drr3 KIirradiated" for 60 rnin at0.0 V (against s.c.e.) in1.0 mol dmW3 KI- 165 f 5- 230 + 18- 195 + 18- 140 + 27- 65 +9- 155 + 5- 155 + 6.5- 230 +3- 150 + 3.5+ 13+ 36+ 28+ 12.5+ 35+ 15+ 31+ 22+ 15HgS darkbrownHgS darkbrownHgS blackHgS blackHgS blackHgS slightlydarkenedHgS dark redHgS blackHgS black" Irradiation is by the xenon lamp uia CuCl, filter.Performed in pH 11.5 0.1 mol dmA3NaN03 (N,-purged, unbuffered with xenon lamp irradiation through CuCl,. Lightintensity w 2.5 x lov2 W cm-2)2596Use of potassium iodide as electrolyte led to rapid darkening of the sulphide andalso to solubilisation. The rate of solubilisation as a function of irradiation timeand concentration of potassium iodide was investigated using stirred suspensions ofthe sulphide.The results are shown in fig. 6(a) and (b). The effect of potassiumiodide upon mercury sulphide electrodes was also investigated. To do this, theelectrodes were irradiated in potassium iodide solution and then removed, washedand dried and their photoresponse measured using 0.1 mol dm-3 sodium nitrate aselectrolyte. The results are shown in table 3. By controlling the irradiation timeand concentration of potassium iodide, electrodes can be prepared which are farmore photoresponsive, as judged by the value of the photocurrent, than the redmercury(r1) sulphide electrodes. The parameters which influence optimisation of thephotoresponse also include the wavelength of the irradiating light, the number ofelectrons passed across the electrolyte-mercury sulphide interface and iodide con-centration.A comparison was made of the effectiveness of halide ions in darkeningthe sulphide and it was found that iodide > bromide > chloride.REACTIONS AT MERCURY(II) SULPHIDE PHOTOANODEBLACKENED RED MERCURY@) SULPHIDE I N ALKALINE0.1 mol dm-3 SODIUM NITRATE SOLUTIONFor these experiments a blackened electrode was prepared in the following way.A red mercury(r1) sulphide electrode was irradiated for 60 min with the xenon lampusing copper(I1) chloride as filter solution and 0.1 mol dm-3 potassium iodide aselectrolyte. The voltage of the mercury(r1) sulphide electrode was held potentio-statically at 0.0 V with respect to a s.c.e.The treatment caused a weight loss in theelectrode. The electrode was washed with distilled water and dried. The photo-response of electrodes prepared in this way was examined in 0.1 mol dm-3 sodiumnitrate solution at pH 11.5. Usually the photovoltage of such electrodes is similaror larger than that of the red mercury(I1) sulphide electrode. The blackenedmercury(r1) sulphide electrodes attain level photovoltage faster than the red mercury(I1)sulphide electrodes. The photocurrents of the blackened electrodes can be up toten times greater than those that are attained with red mercury(I1) sulphide electrodes.Introduction of the darkened electrodes to the alkaline sodium nitrate solution causesno mercury solubilisation (within the limits of detection by atomic absorptionspectroscopy).Plots of the variation in photopotential with pH and of the change inapplied potential required to cause an anodic photocurrent to flow with pH aresimilar to those shown in fig. 2 and 3. However, in both cases the slopes of the linesare -0.03 V pH-l, i.e., half that observed for red mercury(I1) sulphide electrodes.The spectral response of the blackened electrodes differs from that of the redmercury(1r) sulphide electrodes in that it extends beyond 700 nm, there being nopositive photovoltages at sub band-gap wavelengths.Irradiation of the blackened mercury(i1) sulphide electrodes and application ofa biassing potential 3 +0.2 V (against s.c.e.) causes gas evolution at the platinisedplatinum counter electrode. That the gas evolved is hydrogen was shown by thefact that it almost immediately precipitates palladium from dilute aqueouspalladium(I1) chloride solution^.^ Gas evolution from the mercury(I1) sulphideelectrode is very slow.This may be due to the fact that mercury(r1) sulphide adsorbs~ x y g e n . ~ The adsorption of oxygen may account for the fact that irradiation of thedarkened electrodes for long periods leads to a decline in photocurrent. If suchpassivated electrodes are removed from the electrolyte, dried and then put back inthe cell they behave in the normal way after re-establishment of the electrolytR. S . DAVIDSON A N D C. J . WILLSHER 2597junction, i.e., the passivation is not caused by chemical change of the surface of themercury(I1) sulphide.AUTHENTIC B LACK MERCURY ( I I) S ULP HI D E (metf2-C I N NAB AR)I N 0.1 mol dm-3 SODIUM NITRATE SOLUTIONIrradiation of platinum electrodes coated with meta-cinnabar did not generate ane.m.f.Application of a biassing potential of +0.1 V (against s.c.e.) producedminute photoanodic currents. Stirring a suspension of meta-cinnabar (2 g) in500 cm3 of 0.1 mol dm-3 potassium iodide solution led to mercury solubilisation( 3 0 0 ~ g c m - ~ after 5 h). Illumination of the suspension with a 125 W mediumpressure mercury lamp increased the rate of solubilisation (1000 pg ~ m - ~ in 5 h).Solubilisation was observed when sodium nitrate was used as electrolyte. Irradiationof a suspension of the sulphide (2 g) in the electrolyte (0.1 mol dm-3, 500 cm3) for5 h gave a solution containing mercury (200 pg ~ m - ~ ) .DISCUSSIONIt has been shown that platinum electrodes can be covered with pigmentarytitanium dioxide and that illumination of the oxide injects electrons into the platinum.6Electrodes covered with mercury(I1) sulphide in powder form behave in a similarfashion, i.e., the sulphide is acting as an n-type semiconductor.A disadvantage ofelectrodes prepared in this way is the inhomogeneous nature of the covering. Theuncovered regions of platinum give rise to short circuiting and reduce the poweroutput of the cells. This short circuiting made it impossible to perform capacitancemeasurements and so obtain a direct measurement of the flat-band potential anddonor density of the red mercury(r1) sulphide. Another problem with these electrodesis that the contact between the sulphide and platinum may not ohmic.The build up of e.m.f.to a constant value on irradiation and its decline to thedark potential on termination of irradiation followed a similar pattern to thatobserved with titanium dioxide. As with titanium dioxide the build up of photo-currents on commencement of irradiation was dependent on the type of electrolyte.'Consistent results were obtained with sodium nitrate as electrolyte. That this saltdid affect the photoelectrochemistry of mercury(i1) sulphide was shown by the factthat the photoresponse of the electrode was not affected when the concentration ofthe nitrate solution was varied from 10 to 1 xInspection of fig.1 shows that the onset of the anodic photocurrent in 0.1 mol dm-3sodium nitrate solution occurs when the applied voltage is in the range -0.15 to-0.35 V (against s.c.e.). This onset potential is linearly related to the pH of thesolution (fig. 2) as is the photo-e.m.f. of the sulphide electrode. These two parameterschange in a similar way with change in pH. This suggests that illumination of thesulphide electrode under open circuit conditions leads to band flattening and con-sequently the potential required to cause the onset of the anodic photocurrent canbe taken as the flat-band potential.From fig. 3 it can be seen that the dark potential and the potential produced uponillumination vary by x -0.06 V pH-', i.e., the Nernst equation is obeyed.Therelative photopotential is invariant with pH. The fact that the dark potential andpotential produced upon illumination are dependent upon pH to the same extentshows that the conduction and valence bands at the semiconductor-electrolytejunction shift equally to more negative potentials as the electrolyte pH increases.Since the dark potential is a measure of the Fermi level in the dark (against s.c.~.)'the position of the Fermi level at any pH can be potential calculated and put on anmol dm-32598energy scale. The photopotential attained with intense radiation should correspondto the flat-band potential.In order to utilise these values to construct an energy level diagram for redmercury(I1) sulphide the energy of the conduction band was to be determined.Butlerand Ginley have described a method which is based on a knowledge of the solid-stateproperties of the constituent atoms.' Thus the energy of the conduction band ofmercury(1r) sulphide may be calculated as follows.For sulphur : 1st ionisation potential = 10.36 eV and electron dlinity = 2.1 eV.9so,REACTIONS AT MERCURY(II) SULPHIDE PHOTOANODEx(Hg) = +(1.54+ 10.43) = 5.98xq) = 3(2.1+ 10.36) = 6.23a * * X(HgS) = JxfHgS)X(S) = 6*11where x is the electronegativity.level.The undoped Fermi level of mercury(I1) sulphide lies 6.1 1 eV below the vacuumSince the suiphide (red) has a band gap of 2 eV loEA(HgS) = 6.1 1 - +2= 5.11 eVwhere EA is the electron affinity.controlled by the band gap.one uses the equationThe photoaction spectrum (fig.4) confirms that the wavelength response isTo relate the electron affinity value to solution energy levels (redox potentials)E = a-qVwhere E is the energy on the vacuum scale in eV, a is a constant, q is the electroniccharge and V is the potential in Volts in the standard hydrogen electrode (s.h.e.)scale. The constant a has not been definitively fixed and values of -4.48 and-4.73 I 2 have been used. Use of the latter value gives the energy of the conductionlevel (EcB) as +0.37 against s.h.e.Fig. 7 is an energy level diagram drawn up from the values given in table 4.Owing to the uncertainty in the value for the electron affinity for sulphur andthe constant relating the vacuum and s.h.e. scales, these values could vary by kO.3 V.The energy level diagram satisfactorily explains why only a small biassing potentialis necessary to cause the photoassisted electrolysis of water and indicates thatreducing agents with Eo redox negative of the valence band top can be thermo-dynamically oxidised on illumination.TABLE 4.-vALUES OF CONDUCTION BAND BASE (EcB), VALENCE BAND TOP (&IS), FLAT-BANDPOTENTIAL (EFB) AND FERMI LEVEL (&(DARK)) (FOR SEMICONDUCTOR AQUEOUS ELECTROLYTEJUNCTION IN THE DARK) ON VARIOUS SCALESscale ECB EVB EFB(approx) EF(DARK)(approx)vacuum + 5.10 eV +7.10 eV +5.18 eV + 5.38 eVs.h.e., pH = 0 +0.37 V +2.37 V +0.45 V +0.65 Vs.c.e., pH = 7 -0.29 V +1.71 V -0.20 v 0.0 R .S . DAVIDSON AND C. J . WILLSHERtED(cathodic 1E; E; - - - - - -t‘D(anodic12599EC-FIG.7 . 4 ~ ) Postulated band edges at the flat-band condition for red HgS at pH 7 (against s.c.e.)and their relative positions to some solutions redox couples and &(cathodic) and EC(anodic), thethermodynamic potentials for HgS+ 2e + Hg+ S2- and Hg2- + S + 2e + HgS, respectively. 0.15 Vis the energy gap between the semiconductor Fermi level and solution redox level, taken as half waybetweenE(H+/H2) and E ( o ~ / H ~ o ) for an aqueous solution containing no added redox species. (b)Bandbending [0.15 V from (a)] of red HgS. 0.35 V corresponds to the energy needed to raise the conditionband surface edge to the H+ reduction level and 1.15 V is the difference between the 02/H20 levelon the valence band top.Photosolubilisation of HgS, as indicated in table 1, takes place in a number ofelectrolytes, but probably does not occur by anodic oxidation :HgS = Hg2++S+2e.Passage of two electrons should solubilise one mercury atom and z 10l9 electronsneed to be passed to produce a detectable 10 pg ~ m - ~ by atomic absorption spectro-metry.Some electrolytes show [Hg] > 10 pg ~ m - ~ for < l O I 9 electrons passed anda lot of mercury solubilised in suspension where only photochemical and not photo-electrochemical reactions can occur. The “ inert ” sodium nitrate electrolyte withred and blackened sulphide has been employed to pass sufficient current to givedetectable mercury concentrations in solution, but none has been seen. We thereforeconclude that solubilisation is photochemical and any contribution from anodicoxidation is slight, if occurring at all.The relative positions of decompositionlevels l3 with band edges on fig. 7 confirm this, cathodic decomposition is not likelysince ED(cathodic) < Ec, but anodic dissolution is possible as ED(anodic) < E,, butredox couples with Eo < &)(anodic) can compete with the decomposition rea~ti0n.l~Photochemical solubilisation occurs when the sulphide lattice is ruptured by thereaction of a photogenerated positive hole with a reducing agent, but no current ispassed through the lattice. This is the situation for an irradiated suspension.Photoelectrochemical solubilisation (anodic oxidation) is the process of lattice ruptureby passage of electrons through it into an external circuit.The use of potassium ferrocyanide and thiocyanate as electrolyte produces positivephotopotentials.A similar situation arises when platinum electrodes covered withtitanium dioxide are irradiated in potassium iodide, potassium thiocyanate an2600 REACTIONS AT MERCURY(II) SULPHIDE PHOTOANODEpotassium ferrocyanide solution^.^^ It would appear that in all these cases theanions are being oxidised at the semiconductor and the oxidised form of the ions arereduced at the base platinum sites of the semiconductor covered electrode. Thisprocess has been termed " short-circuiting ". The presence of reducible ions in thevicinity of the inhomogeneously covered platinum covered electrode will give riseto a positive photovoltage.The use of potassium iodide as electrolyte caused solubilisation of the redmercury@) sulphide. This was accompanied by the production of high currentsand also a change in the colour of the sulphide to a brown-black colour.The degreeof darkening is dependent upon the duration of the irradiation, the light intensity andthe wavelength of the light. As yet the mechanism of the darkening and the chemicalcomposition of the darkened material are not known. This is a problem which hasintrigued chemists,16 mineralogists and artists.'However, the darkened form of mercuryfrr) sulphide is far more photoreactivethan the red form as shown by its capacity to deliver higher photocurrents. Anothergreat advantage as far as utilisation of the semiconductor for harnessing solar energyis concerned is that it responds to light of A > 700 nm.So far we have not beenable to find the wavelength at which the photoresponse commences. An unexplainedfeature of these blackened electrodes is the effect of changes of pH upon their per-formance. The finding that the variation in photopotential with pH is -0.03 VpH-l is most unusual and suggests different ion adsorption on the blackened formof red mercury(I1) sulphide.The observation that authentic black mercury(I1) sulphide (meta-cinnabar)generates only a slight photoeffect and undergoes considerable photosolubilisationin sodium nitrate electrolyte leads to the conclusion that very little " meta-cinnabar "is formed on darkening. The darkened form is not unstable in a nitrate electrolyteand its nature remains uncertain.X-ray powder photographs and opto-acousticspectra of blackened red mercury(I1) sulphide suggest some " meta-cinnabar " maybe present, although mercury is also solubilised in the iodide-induced darkening, butwe cannot state the exact nature of the blackened form.We thank the S.R.C. for a maintenance grant to C. J. W.R. S. Davidson, R. R. Meek and R. M. Slater, J.C.S. Faraday I, 1979, 75,2507.* R. S. Davidson and C. J. Willsher, British Patent 7913420.R. S. Davidson and C. J. Willsher, Nature, 1979, 278, 238.Comprehensive Inorganic Chemistry, ed. A. F. Trotman-Dickenson (Pergamon Press, Oxford,1973, vol. 1, p, 8.L. I. Grossweiner, J. Phys. Chem., 1955, 59, 742.R. M. Slater, Ph.D. Thesis (University of Leicester, 1975).(a) M. A. Butler and D. S. Ginley, Chem. Phys. Letters, 1977, 47, 319 ; (b) M. A. Butler andD. S. Ginley, J. Electrochem. SOC., 1978, 125, 228.V. I. Vedenyer, L. V. Gurvich, V. N. Kondrat'yev, V. A. Medvedev and Y . L. Frankevich,Band Energies, Ionisation Potentials and Electron Afinities (Edward Arnold, London, 1966).F. Lohmann, 2. Naturforsch. A, 1967, 22, 843.R. Gomer and G. Tryson, J. Chem. Phys., 1977, 66,4413.' S. R. Morrison, The Chemical Physics of Surfaces (Plenum, New York, 1977), p. 269.lo W. H. Strehlow and E. L. Cook, J . Phys. Chem. Ref. Data, 1973, 2, 163.I3 W. M. Latimer, Oxidation Potentials (Prentice Hall, Englewood Cliffs, N.J., 1938), p. 166.I4 (a) H. Gerischer, J . Vac. Sci. Technol., 1978, 15, 1422 ; (b) A. J. Bard and M. S . Wrighton,J. Electrochem. SOC., 1977, 124, 1706 ; (c) H. Gerischer, J. Electroanalyt. Chem., 1977, 82, 133.H. H. Chambers, R. S. Davidson, R. R. Meek and R. M. Slater, J.C.S. Faraday I , 1979, 75,2517R . S . DAVIDSON A N D C . J . WILLSHER 2601l6 C. Brosset, Naturwiss., 1936, 24, 813.l7 W. H. Cropp, Proc. Austral. Inst. Mining Met,, 1923, 52, 259.l8 R. L. Feller, Nat. Gallery of Art, Report and Studies in History of Art (U.S. Govt. PrintingOffice, Washington. D.C., 1967), p. 99.(PAPER 9/1072

ISSN:0300-9599    

DOI:10.1039/F19807602587

出版商:RSC    

年代:1980

数据来源: RSC

摘要:

CORRIGENDUMHornologation of n-Alkanes on Metal FilmsA Novel Aspect of Metal-Carbene ChemistryBy C. O'DONOHOE, J. K. A. CLARKE and J. J. ROONEYJ.C.S. Faraday I, 1980,76, 345-356Correspondence regarding this paper should be addressed to Professor J. J. Rooney atDepartment of Chemistry, The Queens University of Belfast, Stranmillis Road, BelfastBT9 5AG.260

ISSN:0300-9599    

DOI:10.1039/F19807602602

出版商:RSC    

年代:1980

数据来源: RSC