ISSN:0300-9599    

DOI:10.1039/F197369FX001

出版商:RSC    

年代:1973

数据来源: RSC

摘要:

Mechanism of the Steam Reforming of Methane over aCoprecipitated Nickel-Alumina CatalystBY J. R. H. Ross AND M. C. F. STEELSchool of Chemistry, University of Bradford, Bradford BD7 I DPReceived 30 June, 1972The kinetics of the steam reforming of methane over a coprecipitated Ni/A1203 catalyst have beenexamined in the temperature range 773 to 953 K and in the pressure range 0-10 Torr. Examinationof the stoichiometry of the reaction showed that a catalyst freshly reduced in hydrogen at 873 K wasfurther reduced by the reaction mixture. This is taken to imply that reduction of some phase such asNiA1204 was occurring. The rate determining step of the reaction on the fully reduced catalystunder reducing conditions was found to be the rate of adsorption of methane and competition for theadsorption sites by water occurred.The water-gas shift reaction did not proceed appreciably, andthis implies that significant CO and COz adsorption does not occur on the catalyst surface ; this is inagreement with the kinetic results. Experiments involving D,O or D, helped to confirm theseconclusions.The catalysed reaction of methane with steam to produce hydrogen is of consider-able economic importance, as is the steam reforming of higher hydrocarbons. Themethane steam-reforming process is represented by eqn (1) :CH,+H20--+3H,+C0. (1)The CO formed may react with further water by the water-gas shift reaction :CO + H20-+C0, + H2. (2)In the steam reforming of higher hydrocarbons, it is thought that the hydrocarboninitially breaks down to form methane, and that reactions (1) and (2) subsequentlyreach equilibrium.1-6 In industrial use, the catalyst used for these reactions isalmost invariably nickel based; it may be of either the impregnated or the copreci-pitated type, and complex oxide supports are often used in order to reduce the lay-down of carbon and the formation of polymeric material on the surface.There have been several kinetic studies of the steam reforming of methane, eachof which indicates that the rate determining step for the reaction involves methaneadsorption.Akers and Camp,' using an integral flow reactor at atmospheric pres-sure, found that the kinetics of the reaction over a Ni/Kieselguhr catalyst at 911 Kwere first order in methane partial pressure, and independent of water and productpartial pressures.Bodrov et al., using a circulating flow reactor, also at atmosphericpressure, again found first order behaviour with respect to methane partial pressurewith nickel foil * and with two different nickel catalysts.g* lo In all three systems, adependence of the rate on PHI0, PH2 and Pco was observed, although this was greateston the Ni foil; inhibition by hydrogen was important at lower temperatures butdisappeared above 973 K.1°These limited results clearly indicate that the catalyst structure has a marked effecton the kinetics of the reaction. Similarly, it has been shown lq6 that the steamreforming of higher hydrocarbons is dependent on the structure of the catalyst ; e.g.the reaction of butane over nickel on alumina catalysts has very different kinetics1J.R . H. ROSS AND M. C. F. STEEL 11to that over a nickel on urania c a t a l y ~ t . ~ The present study was therefore undertakento examine the effect of a systematic variation of the catalyst composition on thekinetics of the steam reforming of methane. Attention has also been focussed onchanges of the catalyst composition and of the stoichiometry of the reaction withcatalyst use. The present paper reports the results obtained with a coprecipitatedNi/alumina catalyst similar to those used commercially l1 ; the changes in kineticsand stoichiometry of the reaction under a variety of conditions are related to changesoccurring in the catalyst. Subsequent papers will examine changes in the kinetics ofthe reaction when the catalyst formulation is systematically varied.EXPERIMENTALAPPARATUSThe apparatus used was constructed largely of Pyrex ; metal valves were used to isolatethe constant volume reaction system from the pumps and gas-handling line, and the reactionvessels were made of quartz. Analysis was carried out using a low resolution mass spectro-meter, attached to the reaction system by a capillary leak and continuously pumped byseparate pumps.Both the reaction system and the mass spectrometer were bakeable to550 K, and background pressures of < lo-* Torr (1 Ton = 133 N m-,) could be attainedafter bakeout and suitable trapping. Oil diffusion pumps and oil manometers were used inthe system in order to avoid the need for trapping between the reaction system and the gashandling line ; the vapour pressure of the oil at 298 K was - 5 x lo-' Torr.The catalyst was pIaced in one of two matched quartz reaction vessels.Both vesselswere situated in a furnace whose temperature was regulated by a thermocouple and a temper-ature controller. A magnetically operated sliding valve arrangement was used to admit thereaction mixture to either vessel. A magnetically operated centrifugal pump was alsoincorporated to give more efficient circulation of the reaction mixture at higher pressures(> 5 Torr). The total volume of the reaction system, including the reaction vessel, was about4.3 x m3.The whole system was maintained in an oven at 400 K to minimise adsorption of watervapour on the walls and hence to increase the accuracy of the measurement of water partialpressure. Even under these conditions, the response of the mass spectrometer to changes inwater partial pressure was not rapid, and for this reason, the rate of change of water partialpressure was never completely representative of its behaviour in the reaction.Although itwould have been desirable to correct for this, no correction was applied because the effect iscomplex (depending on various factors such as total pressure, water partial pressure, emissioncurrent, etc.). However, it is small and unimportant in the present circumstances, as quotedrates of reaction are based on methane partial pressures.MATERIALSA coprecipitated Ni/alumina catalyst, supplied by Laporte Industries Ltd., containing75 % Ni in its reduced form, was used in this work.This was prepared by a proceduresimilar to that outlined in ref. (1 l), and involved treatment of an aqueous solution of alumin-ium and nickel nitrates with Na2C03, followed by calcining in air at 723 K. The reducedcatalyst contained 0.22 % Na and 0.73 % K. The pelletted catalyst was powdered beforeuse and particles in the range 250-355 pm were used for the experiments reported here.X-ray powder diffraction photographs showed that the catalyst in its unreduced formconsisted mainly of NiO and poorly crystallised y-alumina. The total area of the reducedcatalyst was 175 mZ g-l (N, adsorption at 78 K, B.E.T. method) and the active metal areawas 37 m2 g-' (H2 adsorption at 298 K).Two 0.2 g samples of the catalyst were used andboth gave similar results; one was used mainly for the investigation of the behaviour of afreshly reduced catalyst, while the other was used for the kinetic results on a thoroughlyreduced catalyst, and had been used for preliminary work.CH4, Hl, CO and C 0 2 were supplied in sealed ampoules by the British Oxygen Co12 STEAM REFORMING OF METHANE(Grade X). Freshly deionised water was used after several freezing, pumping and thawingcycles in the gas handling line. D20 was supplied by Prochem Ltd. (99.7 %) and Dz wassupplied by Matheson Inc. (c.P. grade).P I t 0 CE D U R EThe apparatus was calibrated for both reactants and products ; calibrations of the massspectronieter were repeated at frequent intervals.Before each experiment or series of experiments, the catalyst sample was reduced in Hz(-25 Torr) at a temperature of 873 K for several hours.The criterion taken for completereduction was that no water should be produced when the sample was exposed to a freshdose of hydrogen (- 1020 molecules of H2). The reaction system was then pumped, thecatalyst vessel was isolated from the remainder of the system, and a known mixture of reac-tants was admitted to the reaction system and the blank reaction vessel. Complete mixingof the reaction mixture was confirmed and the mass spectrometer calibration was checkedonce a steady state had been achieved. The reaction mixture was then admitted to thereaction vessel containing the catalyst at the desired reaction temperature and the massspectrum was scanned periodically.From these results, graphs showing the partial pressuresof reactants and products as a function of time were constructed.Blank experiments showed that no reaction between CH4 and HzO occurred on the wallsof the quartz reaction vessel or on a sample of y-alumina in the temperature range of theexperiments carried out (773 to 953 K).RESULTSSTOICHIOMETRY OF THE CH4+H20 REACTION OVER A REDUCEDCATALYSTPreliminary work showed that although measurable rates of reaction occurredbelow 773 K, these were not reproducible and depended markedly on the catalystpretreatment. However, above 773 K, reproducible rates were obtained. Under theconditions used, i.e.T = 773 to 953 K and P = 0-10 Torr, the equilibrium in eqn (1)is well to the right, whereas CO and C02 may be expected in roughly equal propor-tions (eqn (2)).Fig. l(a) shows a typical experiment for the reaction of CH4 with H20 at 873 Kon a freshly reduced catalyst ; the reaction mixture consisted of -4.7 x 1019 moleculesor about three molecules of each reactant/ten surface Ni atoms. The rate of disap-pearance of methane after 30 s was - 5 x 1015 molecule s-l. Little water reactedand CO and CO, were produced in comparable proportions. The amounts of COand CO, are compatible with the loss of methane, but not with that of water, and thisindicates that oxygen must be produced from the catalyst. In subsequent experi-ments, carried out without further reduction of the catalyst in hydrogen, the consump-tion of water increased and the proportion of C 0 2 in the reaction products decreasedsteadily. Fig.l(b) shows the results obtained after an extended series of experiments.C02 is no longer a significant product of the reaction, and the stoichiometry of eqn (1)is approximately obeyed. At this stage, 1021 oxygen atoms had been produced fromthe catalyst, which corresponds to - 10 atoms/surface Ni atom. Once established,this stoichiometry of reaction was maintained reproducibly throughout subsequentexperiments, and the results reported below were obtained under these conditions.Fig. l(b) shows that the partial pressure of water does not appear to fall as rapidlyas that of methane; careful consideration of the mass balance of the reactants andproducts shows that this is due to a slow response of the mass spectrometer to changesin water partial pressure.Similar considerations do not apply to the other reactantsand products, and hence the stoichiometry of eqn (1) appliesJ. R. H . ROSS AND M. C. F . STEEL 13100 2 0 0 3 0 0 4 0 0 5 0 0timeis(a)100 2 0 0 3 0 0 420 5 0 0time/s(b)FIG. 1.-Reaction of CH4 with H20 at 873 K over (a) a freshly reduced catalyst and (b) a catalyst forwhich eqn (1) holds. Note that the water partial pressure is subject to some error due to slowresponse of the mass spectrometer.FIG. 2.-Plots (a) of log (rate) against log PCH~ at various values of PH20 and (6) of log (rate) againstlogPH20 at various values of P C H ~ (T = 873 K)14 STEAM REFORMING OF METHANEDEPENDENCE OF THE REACTION RATE ON THE PARTIAL PRESSURES OFREACTANTSIf the kinetics of the reforming reaction adhere to eqn (3),and if we assunie for the present that products do not affect the rate of reaction, thenn and m may be obtained from plots of log (rate) against log PCH4 and log (rate) againstlog PHz0 respectively, with PHLO and PCH4 held constant in turn.The rates used wereobtained after reaction for - 30 s from plots such as those shown in fig. 1(b) ; asdiscussed below, reproducible results were not achieved until that point. Fig. 2(a)and (b) show plots of the data, from which the following values of n and m were found :1.2 = 1.0, 112 = -0.5.The possible error in both values was about 5 0.1.eqn (1) holds throughout the experiment, then eqn (3) becomesIf the initial pressures of methane and water are equal, and the stoichiometry ofand this may be integrated to givePg$4 = const.- kt/2. ( 5 )Fig. 3(a) shows a plot of this relationship for the data of fig. l(b) and also for 14additional experiments carried out under identical conditions at various stagesthroughout the whole experimental programme ; the plot shows satisfactory repro-ducibility from experiment to experiment. All the plots of this type which we haveobtained show a deviation from linearity in the first 30 s of reaction. This deviationdepends largely on reaction conditions, and will be discussed below ; it is because ofthis deviation that the values of 12 and in reported above were obtained from ratesafter - 30 s.DEPENDENCE OF T H E RATE O F REACTION ON THE PARTIAL PRESSURESOF PRODUCTS(a) HYDROGEN.-Initial rate measurements show that the rate of the reaction isapparently retarded by addition of H2 to the reaction mixture.However, when thedata is plotted according to eqn (5), it is seen that the rate is affected only for the first30 s, and thereafter the value of Ic is constant (see fig. 3(b)).(b) CARBON MONOXIDE.-Fig. 3(c) shows similar plots when CO is added to thereaction mixture. Variation of Pco not only affects the initial rate of the reactionbut also causes a slight change in k ; this dependence is given byand is sufficiently small not to affect the general applicability of eqn (3) and hence ofeqn (5).(c) CARBON DIOXIDE.-Fig.3(d) shows the results obtained when C 0 2 was added,and these were similar to those for CO addition ; it was found thatiC cc p;p4 (6)(7) kKp-0.05 co15 J . R . H . ROSS A N D M . C. F. STEEL2.50 . 5 I 1 I I Itime/s(42.510 5 3 100 150 2 0 0 2!time/s(c)time/s@)FIG. 3.-Plots of P$z4 against time for the CH4+ H20 reaction at 872 K for equal initial pressures ofCH4 and H20 of 2.36 Torr showing : (a) reproducibility of the data ; (b) effect of Hz addition : V,4.72; 0, 2.83; A, 0.65; 0, 0.00Torr; (c) effect of CO addition: A, 4.72; 0, 2.36; a, 0.00Torr ; (d) effect of COz addition : V, 4.72 ; 0, 2.36 ; A, 0.65 ; 0, 0.00 Torr.TEMPERATURE DEPENDENCE OF THE CH,+H20 REACTION RATEValues of k (eqn (3)) were determined for temperatures in the range 773 to 953 Kand the results were plotted as log k against 1/T.A value of the activation energy Efor the reaction of 29.0 kJ mol-1 was obtained. Evaluating the pre-exponentialterm, the rate of reaction of methane is given by :where dnldt is expressed in molecules of methane reacted per m2 of surface per secondand the pressures are expressed in Torr (cf. eqn (3))16 STEAM REFORMING OF METHANEREACTION OF METHANE ALONE OVER THE NICKEL CATALYSTWhen a well reacted catalyst, such as that which gives the behaviour shown infig. I@), was exposed to methane alone at a temperature of 873 K, hydrogen wasproduced. The rate of the Ni/CH4 reaction (see fig. 4) was greater than that of theCH4 + H20 reaction.The rate of the carburisation reaction was directly dependenton P&&, as shown by the plot of log PCH4 against time, also shown in fig. 4. Incorpor-ation of carbon from the methane was not restricted to one dose (2.4 x l O I 9 molecules),but continued with a gradual decrease of rate for many doses, equivalent to -30monolayers of “ carbide ”.I 1FIG. 4.-Comparison of the rate of reaction ofCH4 with the reduced catalyst (0) with that ofthe CH4 + H20 reaction (a). The first orderdependence of the CH4 reaction is also shown(A) (T = 873 K).0 . 0 00 100 2 0 0 3 0 0 400timelsREACTION OF WATER WITH A “CARBIDED” CATALYSTCO and H2 were both produced when water vapour was reacted with a “ car-bided ” catalyst at 873 K but no methane was observed; nor was there any evidencefor the loss of oxygen to the catalyst.The reaction may therefore be representedby eqn (8) :“ carbide ” + H,O-+Ni + H2 + CO.Fig. 5 shows that the rate of reaction of water with the “ carbide ” was considerablygreater than that of the CH,+H,O reaction on the “ carbided ” catalyst and wasalso more rapid than the latter reaction on a reduced catalyst.(8)KINETICS OF THE CH,+H,O REACTION ON THE “ CARBIDED ” CATALYSTThe stoichiometry of the reforming reaction on the “ carbided ” catalyst wassimilar to that shown in fig. l(b).molecule m-2 s-l at 873 K, as opposed to - 5 x 1015 molecule s-l under similarconditions for the reduced catalyst (see fig. (5)). The following values of n and m(eqn (3)) were obtained :However, the rate was appreciably less (-3 xn = 1.0 m = 1.0.As before, the values of n and m were considered to be accurate to about kO.1J .R. H. ROSS AND M. C . F . STEEL 17EXCHANGE EXPERIMENTSA series of experiments was carried out with CH4+D20 and CH4+H20+D2mixtures, in order to examine any exchange that might occur in the methane con-current with the steam reforming reaction.Negligible exchange of the methane was observed during its reaction with D20,the reforming reaction being at least ten times faster than any exchange reaction.Both D2 and H2 were formed and complete equilibration giving HD was also ob-served. The unreacted D20 at any stage of the reaction was also fully in equilibriumwith product H2 to form HDO and H20.Similar results were obtained in the CH4 + H20 + D2 experiments : the exchangeof the methane was negligible, but complete equilibration of the H20 and D2 occurred.Complete exchange of H20+D2 mixtures was also found on the blank reactionvessel at 873 K, but the exchange occurred more slowly.2.5FIG.5.4omparison of the rate of the reactionof HzO with a carbided catalyst (0) with thoseof the CH4+H20 reaction on a reduced (a)and a carbided surface (V) (T= 873K). 1.5 -Note that although the water partial pressures --are subject to some error (see experimental s e ~tion), the comparison is still valid, 22.01.0 -Oo5 t0.00 100 2 0 0 3 0 0 4 0 0 5 0 0timelsDISCUSSIONCHANGES I N THE CATALYST COMPOSITION AND REACTION STOICHIOMETRYNickel oxide supported on alumina is notoriously difficult to reduce,l particularlyin the coprecipitated form l3 ; unsupported NiO, however, can be completelyreduced under carefully controlled conditions.l4 After using the criterion for com-plete reduction mentioned above, we still find the equivalent of N 10 monolayers ofoxygen produced from the catalyst during the CH4+H20 reaction, as evidenced bythe production of CO and C02 without loss of water (fig. l(a)). The standard freeenergy change for the reduction of nickel oxide by hydrogen :is -34.81 kJ mol-1 at 873 K, and hence the equilibrium position is well over to theright-hand side. In addition, the reduction of nickel oxide by CO may occur :NiO+H,-+Ni+H20 (9)NiO+CO-+Ni+CO, (10)as the value of AG& for this reaction is -42.05 kJ mol-l18 STEAM REFORMING OF METHANEHowever, we must also consider the possibility of the formation of nickel alum-inate, NiA1204, when the catalyst is calcined in air at 723 K during preparation,because its presence has been reported at temperatures as low as 673 K in similarl6 For the reduction of NiA1204 by hydrogen :NiA1204+H2-+Ni+A1203 +H20 (1 1)NiA1204+CO+Ni+A1203 +C02 (12)AG0873 is - 15.73 kJ mol-l, and for the reduction by CO :AGg73 is -22.97 kJ mol-l.These values are calculated from the data of Lenev.I7The lower values of AG& for the reduction of nickel aluminate compared to nickeloxide show that the presence of the spinel or similar oxide phases makes the reductionof the catalyst less energetically feasible, but at the same time does not precludecomplete reduction by either hydrogen or carbon monoxide.It is of interest to notethat Delgass has recently obtained evidence from photoelectron spectroscopy datafor the presence of an intermediate compound in the NiOjalumina system, althoughits composition is not specified.The values of AGS73 given for the reactions shown in eqn (9)-(12) apply to stoichio-metric oxides, and may be smaller if non-stoichiometric oxide phases are present inthe region of the surface. Hence, it is probable that the last stages of reduction willnot occur when the reducing agent is hydrogen, as in the initial reduction of thecatalyst, but that when CO is formed during the CH4+H,0 reaction, it is capable ofbringing about the complete reduction of the surface layers.The thermodynamic data also afford an explanation of the initial deviation fromlinearity of the P;g4 against time plots (fig.3). The spontaneous approach to equi-librium of the reactions given in eqn (9) and (12) is governed by :(13) PHlo' H 2AG873 = AG,",,+RT ln-.When AGg73 is negative, the reactions will proceed in the directions shown, but ifAG873 is positive, the reverse reaction will take place. NiO (eqn (9)) is reducedwhen PH20/PH2 < 117 but reduction of NiA1204 will not occur until PkIzo/PHz < 8.6.Until the hydrogen partial pressure is such that these ratios are achieved, oxidationwill occur, and so the initial stages in all the CH,+H,O reactions take place onpartially oxidised catalyst.In studies of the steam reforming of hydrocarbons, it has generally been as-sumed l-l0 that the water-gas shift reaction (eqn (2)) is at equilibrium. The decreasein the production of C02 as reduction of the catalyst proceeds indicates that this maynot be the case in the present work.Fig. 6 shows typical results for the disappearanceof methane as a function of time together with the calculated equilibrium values ofPco, Pco, and PHlo. Pco2 increases initially and then falls again as PHz builds up.Such behaviour was in no case observed in the present work (see figs l(a) and (b)).Initially, COz is formed by interaction of CO with the surface oxide, but Pcoz does notfall as the hydrogen partial pressure increases ; in later experiments, CO productionis predominant.We must therefore conclude that the shift reaction (eqn (2)) doesnot occur to an appreciable extent in the presence of methane. This implies that thesurface species resulting from the adsorption of methane are more strongly bondedto the surface than species arising from CO and C 0 2 adsorption, and this is in agree-ment with the kinetic results (see below).No evidence for loss of carbon to the catalyst was found during the CH, + H20 re-action. On the other hand, when methane was reacted with the catalyst, incorporatioJ . R . H. ROSS AND M. C. F . STEEL 19of carbon occurred, and this uptake of carbon was not limited to the surfacelayer.Nickel carbide (Ni3C) is known to be produced on heating nickel metal in thepresence of carbon containing compounds (e.g.CO, CH4),19 but it decomposesrapidly above 703 Kin vacuum to Ni and graphite.20 Galwey has shown 21 that Ni3Creacts with water at temperatures between 500 and 683 K to give COz, H2 and smalleramounts of CH4 ; the quantities of product formed decrease with increasing temper-ature and this may be due to slower reaction of the graphite formed from the thermaldecomposition of the Ni3C. Galwey's postulate that C 0 2 may oxidise the metal toNiO seems unlikely, however, in view of the thermodynamic data given above. Themain products of the reaction of the " carbide '' with water in the present work areCO and H2 and the reaction is rapid, which would be unlikely if graphite wereinvolved.The relative rates of the " carbide "/H20 reaction and the CH, + H20reaction on the reduced catalyst (fig. 5 ) explain why carbon deposition does not occurduring the reforming reaction, because, even if carbon atoms are formed, they areimmediately removed by reaction with water vapour.timeisFIG. 6.-Calculated distribution of products for a typical rate of disappearance of CH4 at 873 K,assuming thermodynamic equilibrium.REACTION MECHANISMIn agreement with previous investigations of the methane steam-reforming reac-tion, the rate of reaction was found to be first order with respect to PCH4. This impliesthat the rate determining step is the dissociative adsorption of methane. Thedependence of the rate on implies that the water competes with the methanefor the active catalytic sites.This situation is shown in the scheme (see opposite).The rate of formation of CH3 surface species determines the rate of reaction (step 1).Reversible dissociative adsorption of water (steps 2 and 2') occurs competitively onthe same sites. Once surface CH3 species are formed, they may break down furtherto form CH2, CH or C surface entities which may form oxygenated surface species byinteraction with OH groups. These break down to give CO(g); if adsorbed C20 STEAM REFORMING OF METHANEgroups existed, equilibrium amounts of C 0 2 would be formed. Readsorption of COhardly occurs.Methane alone reacts with the catalyst more rapidly than does the CH, +H,Omixture, and this is further evidence that step 1 is inhibited by water.Also, in theabsence of water, step l a proceeds to form surface carbon atoms which becomeincorporated in the lattice. When the reforming reaction is carried out on thecarbided catalyst, the rate is directly proportional to both the water and methanepressures, and this implies that H 2 0 is reacting directly with carbon from the lattice,while CH4 is replenishing this carbon. The sites involved are different to those on thereduced catalyst, and the reaction is slower.SCHEMEThe exchange experiments reported above were carried out to help confirm thereaction scheme proposed above. For example, if any step other than 1 was ratedetermining and step 1 was reversible, then the mode of the exchange within theunreacted methane molecules would show which step was important.The lack ofexchange with either D,O or D2 confirms that step 1 must be rate determining, andthat the reverse reaction does not occur. Complete exchange of the hydrogen andwater species shows that steps 2 and 2’ and 3 and 3’ are rapid. It is of interest to notethat Morikawa et aLZ2 studying the CH4+D,0 reaction over a Ni/Kieselguhrcatalyst at 500 K, found slow exchange concurrent with an equally slow reformingreaction in which H2 and C02 were produced. However, their highest temperatureof reduction in pure hydrogen was 723 K, and our results suggest that this may not be asufficiently rigorous treatment to produce a catalyst of high reforming activity.The authors thank Professor M. W. Roberts for his interest and encouragement,and the General Chemicals Division of Laporte Industries Ltd. for generous support.M. C. F. S. thanks the S.R.C. for an award under the C.A.P.S. scheme.C. H. Riesz, H. A. Dirksen and W. J. Pleticka, Inst. Gas Tech. Res. Bull., 1952, 20.M. C. F. Rogers and W. M. Crooks, J. Appl. Chem., 1966, 16,133.K. S. M. Bhatta and G. M. Dixon, Trans. Faraday SOC., 1967, 63,2217.K. S. M. Bhatta and G. M. Dixon, Ind. Eng. Chem. Prod. Res. Deu., 1969, 8, 324.T. R. Phillips, T. A. Yarwood, J. Mulhall and G. E. Turner, J . Catalysis, 1970, 17, 28.C. R. Schnell, J. Chern. SOC. B, 1970, 158. ’ W. W. Akers and D. P. Camp, Amer. Inst. Chem. Eng. J., 1955, 1,471. * I. M. Bodrov, L. 0. Apel’baum and M. I. Temkin, Kinetika i Kataliz, 1964,5, 696.I. M. Bodrov, L, 0. Apel’baum and M. I. Temkin, Kinetika i Kataliz, 1967, 8, 821.lo I. M. Bodrov, L. 0. Apel’baum and M. I. Temkin, Kinetika i Kataliz, 1968, 9, 1065.l1 Brit. Pat. 969,637 1961.l2 D. Reinen and P. W. Selwood, J. Catalysis, 1963, 2, 109.l3 V. C. F. Holm and A. Clark, J. Catalysis, 1968, 11, 305.l4 M. W. Roberts and K. W. Sykes, Trans. Faraday SOC., 1958, 54, 548.l5 A. M. Rubinshtein, V. M. Akimov and L. D. Kretalova, Izuest. Akad. Nauk S.S.S.R., Otdel.khim. Nauk, 1958, 929J . R. H . ROSS AND M . C . F. STEEL 21l 6 M. Lo Jacono, M. Schiavello and A. Cimino, J. Phyx. Chem., 1971,75,1044.l7 L. M. Lenev and I. A. Novokhatskii, Zhur. neorg. Khim, 1965, 10,2400.W. N. Delgass, T. R. Hughes and C. S. Fadley, Catalysis Reu., 1971, 4, 179.l9 L. J. E. Hofer, E. M. Cohn and W. C. Peebles, J. Phys. Colloid Chem., 1950,54, 1161.2o S. Nagakura, J. Phys. SOC. Japan, 1957,12,482.'* K. Morikawa, W. S. Benedict and H. S. Taylor, J. Amer. Chem. SOC., 1936,58, 1445.A. K. Galway, J. Catalysis, 1963, 2, 176

ISSN:0300-9599    

DOI:10.1039/F19736900010

出版商:RSC    

年代:1973

数据来源: RSC

摘要:

Adsorption of Carbon Monoxide by Zeolite Y Exchanged withDifferent CationsB Y T. A. EGERTON~ AND F. s. STONE*$School of Chemistry, University of Bristol, Bristol BS8 ITSReceizied 5th July, 1972Adsorption of carbon monoxide has been studied on Zeolite Y in which Na+ ions have beenpartially exchanged for Zn2+, Mnz+, CoZ+, Ni2+, Cu2+, Ba2+, UO:' and Ce3+ ions. The divalentions all produced sites for specific CO adsorption, even at low degrees of exchange. Thus, unlikeCa2+ studied in earlier work, these ions do not have a total preference for internal sites in the prismsor sodalite units inaccessible to CO. At approximately 30 % exchange the affinity of divalent ionsfor internal sites decreases in the order Ca2+ > Ni2+, Mn2+, UOg+ > Cu2+ > Zn2+. However,even on the Zn-exchanged Y , only about 1 divalent ion in 10 acts as a site for specific adsorption ofCO.Very little CO was specifically adsorbed on Ce-exchanged Y, even at high degrees of exchange.This confirms that cerium ions favour internal sites.With ZnY, MnY, BaY and CeY, CO adsorption was rapid and reversible. Isosteric heats wereevaluated and, except for CeY, were substantially greater than on unexchanged Nay, confirmingthe presence of specific adsorption. The heats correlate with the electrostatic field strengths of thecations and with the shifts of the C-0 stretching frequency on adsorption. With NiY and CuY,adsorption was slow and is thought to reflect a gradual increase in the number of adsorption sitescaused by adsorbate-induced migration of divalent cations.A number of the systems studied in this work have not previously been investigated from thestandpoint of cation location.However, where comparisons are possible, the present results derivedfrom CO adsorption are shown to be in good agreement with those from other methods.The principal methods which have so far been used to study the location and co-ordination taken up by the charge-balancing cations in ion-exchanged syntheticfaujasites are the more obvious physical ones. Thus X-ray diffraction,l* electronre~onance,~ infra-red spectr~scopy,~ reflectance spectr~scopy,~~ Mossbauer spectro-scopy ' and magnetic susceptibility * are among the methods which have been appliedin individual cases. These methods are entirely sufficient in principle, and for themost part experimentally straightforward, but the interpretation of the results andthe establishment of firm conclusions is frequently difficult.The aim of the presentwork is to show that physico-chemical studies of specific adsorption can also be usedto give information on the siting of cations. The method is not as powerful as someof the physical ones, but it has the advantage of being applicable to a wide variety ofions. Moreover, it leads us closer to the chemical context of the problem, which isof growing importance now that many zeolites are being used as catalysts. Quiteapart from the matter of site occupation, there is much to be gained by studying thechemical interactions which occur between zeolite cations and the reactive gases ofinterest in catalysis.Carbon monoxide is an attractive choice for such work.The CO molecule ist present address : Department of Chemistry, Makerere University, P.O. Box 16020, Kampala,$ present address : School of Chemistry and Chemical Engineering, University of Bath, BathUganda.BA2 7AY.2T. A . EGERTON AND F. S . STONE 23small enough to enter the supercages of faujasite, but too large to enter the sodaliteunits or hexagonal prisms. The molecule has an asymmetric charge distribution andis easily polarized ; it is therefore sensitive to the strong electrostatic fields surroundingcations. It also interacts very specifically with transition metal ions, as a result ofits ability to act both as a weak donor and as a n-acceptor, and it is well known toform chemisorption bonds on transition metal oxides.A few years ago Angell and Schaffer showed by infra-red spectroscopy that COmolecules are indeed adsorbed close to multivalent cations in faujasite supercages, andat the same time Rabo, Angell and co-workers illustrated some of the potentialities ofanalyzing the cation-specific adsorption of CO on zeolites. We have since reporteda detailed study of calcium-exchanged Zeolite Y using CO as a probe molecule,''and in this paper we present results and conclusions on NaY exchanged with eightother ions, including several transition metal ions.Experiments on the influence ofoutgassing conditions have also formed part of the work, since there is increasingevidence to suggest that small amounts of residual water can exert a dominant effecton cation location.EXPERIMENTALMATERIALSAll samples used in the research were prepared from the same batch of synthetic sodiumfaujasite (Nay) which was kindly supplied by the Linde Division of the Union CarbideCorporation (SK-40, Lot No.51-31). The composition was in good agreement with theformula Na, 6(A102)5 6(Si02)13 6,264 H20.Aqueous solutions containing suspended Nay became alkaline on standing, so the solidwas washed with a sodium acetate+acetic acid buffer (pH 5) before being used in an ionexchange reaction. Small amounts of this buffer were also added to the first batch of asolution used in any ion exchange reaction. If these precautions are not observed, the highpH developed within the zeolite pores can lead to the incorporation of a hydroxylated formof some ions. With manganese, for instance, absence of buffer leads to the number ofequivalents of Mn2+ entering the zeolite exceeding the number of Na+ equivalents replaced,suggestive of Mn(OH)+ formation, and over a period of months samples change in colourfrom white to brown.Individual details of the ion exchange procedure employed in the preparation of samplesare summarized in table 1.A suspension of 5 g of pretreated NaY in 200 cm3 of theappropriate solution was stirred for 2 h. When necessary the sodium content was loweredfurther by repeated exchange. The zeolite was then filtered out, thoroughly washed anddried overnight at 70°C. Finally, it was stored over saturated ammonium chloride solutionfor at least three weeks before use.The water content of a fully hydrated ion-exchanged sample was determined within 0.2 %by dehydrating it at 400°C on a McBain balance.The extent of ion exchange was measuredwithin 1 % by flame photometric analysis for residual sodium. With MnY the sodiumanalysis was supplemented by a volumetric (EDTA) determination of manganese, and withNiY the nickel was analysed gravimetrically as the dimethylglyoximate. In both cases thenumber of equivalents of sodium lost equalled the number of equivalents of transition metaIion gained.In certain cases ion exchange can lead to collapse of a zeolite framework, and in additionsome exchanged zeolites which are stable in the hydrated form decompose when outgassed.Accordingly, selected samples were studied by X-ray diffraction both before and after use inadsorption measurements.Diffractograms were measured with a Philips PW 1051 diffracto-meter and Cu-K, radiation, using a sample rotator to improve reproducibility. No X-rayevidence of structural breakdown was observed. Low temperature (78 K) adsorptionisotherms of argon (c = 0.138 nm2) andlor nitrogen (a = 0.162 nm2) were also measure24on some samples. Surface areas (monolayer equivalent areas) were evaluated using bothpoint B l3 and Langmuir plot l4 methods, with the results shown in table 2. Only in thecase of MnY-75 * did the adsorption measurements suggest that substantial structuralbreakdown had occurred.co ADSORPTION ON ZEOLITE YTABLE 1 .-PREPARATION OF ION-EXCHANGED SAMPLEScationBa2fCe3+CO2+cu2+Mn2+Ni2+uo;+Zn2+ionicrlnm ion exchange solution t exchange/ "Cradius 12 temp.of0.134 0.05 M BaCI,0.107 0.005 M CeC130.072 0.1 M C0(N03)20.072 0.1 M C~(N03)20.080 0.05 M MnC120.069 0.005 M NiS042525251 0010010010010025252525252525250.05 M U02(CH3C00)2 100 u-0= 0.1710.074 0.02 M ZnC12 100100% degreeof exchange8417180803242872231407520465866333162t The concentrations are average values.TABLE 2.-sURFACE AREAS, AS DETERMINED BY ARGON AND/OR NITROGEN ADSORPTION AT 78 Ksurface area $argon adsorption NZ adsorptionoutgassing Langmuir Langmuirsample temp./"C point B plot point B plotNayCeY-40CeY-40Cey-71CeY-80CeY-80COY-80CUY-87MnY-75NiY-57ZnY-3 1ZnY-31ZnY-3 1ZnY-763503504803503 5080035035035035034057076035063564063060069070060559063562561061 5675 680 70065064061070569563 5 710 755605570 575648 680655645620635 625 640$ Areas are expressed as m2 g-' of hydrated zeolite.* The symbol preceding Y denotes the ion introduced by ion exchange, and the number denotesthe percentage of Na+ ions in the original Nay which have been replacedT. A .EGERTON AND F. S. STONE 25ADSORPTION MEASUREMENTSThe adsorption measurements were made volumetrically, as previously described. OGreat care was taken during the initial stages of dehydration, and only after many hourspumping was the outgassing temperature raised above 120°C.Except where otherwisestated, the samples were outgassed overnight at 350°C prior to the determination of eachadsorption isotherm.RESULTSThe results are presented in three sections. Section A contains the results ob-tained with ZnY, MnY and Bay, and includes determinations of heats of adsorption.These experiments complement the studies on NaY and CaY already reported. loIn section B we compare the behaviour of NiY, COY and CuY, and illustrate somenew features not found with the cations concerned in section A. In section C weplace CeY; in contrast to the ions in sections A and B, the new characteristic ofZeolite Y exchanged with cerium ions is that the solid shows little evidence of specificCO adsorption.A. CO ADSORPTION ON ZnY, MnY AND BaYFor these zeolites, the adsorption was rapid and equilibrium was always attainedwithin 30 min.The adsorption was reversible without hysteresis and isosteric heatswere calculated in the usual way from plots of (In P)e against 1/T.ZnYThe adsorption of CO on ZnY-31 and ZnY-62 was measured at five differenttemperatures after outgassing at 370°C. The set of adsorption isotherms for ZnY-62is shown in fig. 1, and heats of adsorption for both samples are given in table 3.At 0.6 em3 adsorbed per gram the heat of adsorption on ZnY-31 is 47 kJ mol-1 andthat on ZnY-62 is 52 kJ mol-1 ; both heats considerably exceed the value of 25kJ mol-I for CO adsorption on NaY at a coverage of 0.10 cm3 g-l.The similarityFIG. 1.-Adsorption of CO on ZnY-62 outgassedat 370°C: a, 0°C; e, 22.2"C; 9, 32.2"C;0, 40.9"C; c), 51.7"C.3I M-0 0.5 I .opressue/kN m-26 co ADSORPTION ON ZEOLITE Yin the heats for the two ZnY samples suggests that adsorption occurs on similar sites.Barrer and co-workers l4 have demonstrated that high heats of adsorption of polarmolecules on alkaline earth zeolites result from the strong electrostatic fields associ-ated with the divalent cations; clearly the high heats on ZnY have a similar cause.Previous work on CaY lo has shown that one adsorbed CO molecule is associatedwith each exposed divalent cation. If it is assumed that adsorption on ZnY proceedsin a similar manner it is possible to estimate the number of exposed zinc ions inZnY-31 and ZnY-62. To a first approximation the amount of specific adsorptionon the zinc cations may be calculated by subtracting the CO adsorption on NaY at1 kN m-2 from the corresponding adsorption on ZnY.Isotherms for CO adsorptionat 0°C on NaY and ZnY-31 are shown in fig. 2.TABLE 3.-HEATS OF ADSORPTION OF CO ON ZnY-3I AND ZnY-62heat of adsorption/kJ mol-1ZnY-3 1 ZnY-62coveragelcm3 g-1 8.7 cation1u.c. 17.4 cationlux.0.2 51 .O_+ 2.0 -0.4 48.0+ 2.0 -0.6 46.51 2.0 51.5k4.00.8 43.5k2.0 52.0+ 4.01 .o 39.0f 3.0 51.5+ 2.01.41.82.02.552.5+ 2.0I 5 1 .O_+ 2.0- 44.0+ 3.0- 36.5_+ 3.0-Ic.5 ! .C Ipressure/kN n r 2FIG. 2.-Comparison of 0°C isotherms for:a, NaY ; 0, Cay-36 ; 9, U02Y-33 ; 0 , MnY-31 ; 8, CuY-32; 0 , ZnY-31.For ZnY-31 the amount of specific adsorption calculated in this way is 1.04 cm3 g-l.This corresponds to 0.81 CO molecules per unit cell (u.c.) and hence implies thatthere are 0.81 exposed Zn ions per U.C.on ZnY-31. Altogether there are 8.65 Znions per U.C. of ZnY-31 and if they were all accessible they would be able to adsorb11.2 cm3 g-l of CO. A random distribution of Zn ions would correspond to aspecific adsorption of 4.6 cm3 g-l of CO.* It follows that the majority of the zinc* This is based on a distribution between 16 type I, 32 type I’ and 32 type IT sites, with only oneof a given trio of 11, IIo and 11’ sites occupied at a given timeT. A . EGERTON AND F. S .STONE 27ions are in sites inaccessible to CO. We conclude that they occupy the internal(I, 1’, or possibly 11’ lo) sites in the zeolite structure. On ZnY-62 the amount ofspecific adsorption was 2.73 cm3 g-l, which corresponds to 2.12 exposed Zn ions perU.C. compared with a total of 17.4 per U.C. in the zeolite. Again, the majority of thezinc ions are evidently inaccessible to CO.Fig. 2 affords a comparison of the 0°C isotherms for CO on ZnY-31 and Cay-36.In low-exchanged Cay, the preference of divalent ions for inaccessible sites is virtuallyabsolute.lo It is apparent that, although a high proportion of the Zn2+ ions inZnY-31 is in internal sites, the zeolite has considerably more accessible divalent ionsthan the corresponding calcium zeolite.A further difference between low-exchangedsamples of CaY and ZnY is their response to changes in outgassing temperatures.Increased outgassing temperature led to little change in the amount of CO adsorbedon low-exchanged CaY,l0 but with ZnY-31 outgassing at 570°C instead of 370°Cled to a significant increase in the amount of CO which could be subsequentlyadsorbed (fig. 3). When the isotherms in fig. 3 are corrected for the component ofFIG. 3.-The dependence of the 0°C adsorption onZnY-31 on the outgassing temperature. Out-gassing at : Q, 370°C ; 0 , 405°C ; 0, 505°C;@, 570°C ; 0, 750°C.5non-specific adsorption, it is found that the increased outgassing temperature hasled to an increase in the limiting value of the specific adsorption ; i.e.there are moreadsorption sites.Raising the outgassing temperature probably results in the removal of smallquantities of chemisorbed water from Zn ions in internal sites. There is no room forwater to be chernisorbed on Zn ions in the hexagonal prisms (site I), so we concludethat the inaccessible Zn ions in ZnY-31 are in the sodalite units. On removingchemisorbed water, which is believed to stabilize cations in the I’ and 11’ positions,the stability of ions within the sodalite units will be lowered and this could causethem to move to other sites, such as I1 or IIo, where they are accessible to CO. Itfollows that at least some of the ions in low-exchanged ZnY are in different sitesfrom those occupied by the calcium ions in low-exchanged Cay.On raising the outgassing temperature to 750°C the 0°C adsorption isotherm wasdepressed (fig.3), although separate surface area determinations (table 2) indicatedthat there was no marked loss of crystallinity. There may be an increased tendenc28 co ADSORPTION ON ZEOLITE Yfor the zinc ions to occupy the empty type I sites in the hexagonal prisms as the lasttraces of water are removed.MnYIsotherms were measured over a range of temperatures for MnY-22, MnY-31,MnY-40 and MnY-75. The 0°C isotherm for MnY-31 is compared with the 0°Cisotherms for Nay, Cay-36 and ZnY-31 in fig. 2. In fig. 4 is shown the set of adsorp-tion isotherms for MnY-22. Adsorption on all four samples was reversible, and forMnY-22, MnY-31 and MnY-40 the plots of (log P)o against 1 /T were linear.Themeasured heats are listed in table 4.FIG. 4.-CO adsorption on MnY-22 outgassed at350°C: 0 , 0°C; 0, 19.3"C; 0, 29.4"C; 0 ,42.2"C.c.5 1.0 1.5pressure/kN m-*TABLE 4.-HEATS OF ADSORPTION OF co ON MnYheat of adsorption */kJ mol-1coverage/cm' g-1 MnY-22 MnY-3 1 MllY-40 MnY-750.05 46.5 48.5 50.0 49.5 'f0.10 42.5 45.0 51.00.25 42.5 43.50.40 31.0 35.00.50 32.0* k 2.5 kJ moI-l. 7 This value was calculated from the best straight line drawn through thepoints between 103K/T = 2.36 and 3.18.Compared with Nay and CaY -36 (fig. 2) the appreciable curvature of the adsorp-tion isotherms and the increased heat of adsorption on low-exchanged MnY showthat Mn2+ is not like Ca2+ in having an absolute preference for sites inaccessible to CO.Mn2+ in Zeolite Y has a behaviour intermediate between Ca2+ and Zn2+.The specificadsorption on MnY-22, calculated in the same manner as that for ZnY-31, is 0.15cm3 g-l. This implies that 0.12 of the 6.1 Mn2+ ions per U.C. are accessible to CO :2.4 would be accessible if the manganous ions were randomly distributed. Thus thevast majority of Mn ions in MnY-22 are not available as adsorption sites for CO.Fig. 5 shows the CO adsorption on MnY as a function of Mn content. ThP r . A . EGERTON AND F . s. STONE 292.8FIG.as a- '*1 a5.-Adsorption of CO at 0°C and 100 N m-2 3function of thepegree of exchange : a, CeY ; 20 , MnY; 0, NiY; 0 , CuY. W9 I i-0 2 0 40 6 0 to% exchangeamount of adsorption increases rapidly beyond 50-60 % exchange, which stronglysuggests that a group of inaccessible sites becomes filled at this degree of exchangeand that further Mn ions which are introduced occupy sites accessible to CO.NowMnY-57 corresponds to a zeolite containing 16 Mn ions per U.C. Hence this resultprovides good evidence that the first Mn ions exchanged into the structure have astrong, but not absolute, affinity for a group of about 16 internal cation sites. Thesesites are probably those in the prisms (site I), of which there are 16 per unit cell.For MnY-75 the (log P)e against l / T plot showed appreciable curvature (fig. 6)103 KITand the value of the adsorption heat recorded in table 4 is only approximate. Thelow surface area of this sample (see table 2) indicated that some structural breakdownhad occurred.It i s possible that the curved heat plot is a consequence of the partialcollapse of the framework, although experiments in which the framework of MnY-430 co ADSORPTION ON ZEOLITE Ywas deliberately partially destroyed led neither to curved heat plots nor to a changein the heat of adsorpfion (fig. 6).BaY AND U02YBay-8 (2Ba2+ per unit cell) was outgassed at 350°C and CO adsorption wasmeasured at three different temperatures (fig. 7). The 0°C isotherm was thenFIG. 7.-Adsorption of CO on Bay-8 and U02Y-33 : e, 29.2"C on BaY outgassed at 350°C ; (>,17.9"C on BaY outgassed at 350°C ; 9, 0°C onBaY outgassed at 350°C; 8, 0°C on BaY out-gassed at 425°C ; 0, 0°C on U02Y outgassed at350°C.pressure/kN m-2redetermined after outgassing the zeolite at 450°C ; the higher outgassing temperaturecaused the adsorption to increase.The curved adsorption isotherms and the heatof adsorption of 36 kJ mol-1 (11 kJ mol-1 greater than on Nay) show that eventwo barium ions per unit cell affect the adsorption properties of Zeolite Y. How-ever, the specific adsorption is only 14 % of the amount which would be expected ifall the barium ions acted as adsorption sites. Thus barium behaves in a similarmanner to manganese and zinc.Adsorption on U02Y-33 was measured at 0°C only. The isotherm (fig. 7) showsthat the uranyl ion behaves like manganese and zinc also.B. CO ADSORPTION ON COY, NiY AND CuYThe isotherms for these three zeolites were markedly rectangular with a pro-nounced knee.For CuY and COY only 0°C isotherms were measured. For NiYthere was a gradual decrease in adsorption when the temperature was raised from0°C to 75"C, but as the isotherms were not reversible isosteric heats of adsorptioncould not be evaluated.COYAdsorption on COY-80 was measured at pressures between 0 and 150 N m-?The extent of specific adsorption was 1.5 cm3 g-l. If COY behaved as Cay, with allcations beyond the first 16 occupying accessible sites,1° there would be six exposedcobalt ions per unit cell and the specific adsorption would be - 8.2 cm3 g-l. Thelow adsorption shows that the cobalt ions have a preference for sites within the sodalit31cages, where they are hidden from CO molecules. The presence of cobalt ions intype 11' sites will help to compensate the negative charge on nearby oxygen windows.T .A. EGERTON AND F. S. STONENiY350°C, and on NiY-46 outgassed at 500°C.Adsorption was measured on NiY-20, NiY-46, NiY-57 and NiY-66 outgassed atRepresentative isotherms are shown in fig. 8. Clearly the initial steep sections ofFIG. &-Adsorption of CO on NiY: e, 0°Cadsorption on NiY-58; 0, -42°C on NiY-58;6,O"C on NiY-46 ; 9,O"C on NiY-46 outgassedat 500°C ; 0 , 0°C on NiY-20 ; (>, readsorptionat 0°C on NiY-20 after pumping overnight at 25°C.5the curves are associated with adsorption on the Ni2+ ions. Between 200 and 1500N the slopes of the isotherms are similar to the corresponding NaY plots. Thisindicates that tlie adsorption at high pressure is occurring on the aluininosilicateframework and on the residual sodium ions.The adsorption represented by boththe initial and final straight sections of the isotherms was rapid (equilibrium wasalways attained within 30 min) but that associated with the " knee'' was muchslower and in some cases equilibrium was established only after 36 h.The steep initial section of the isotherms suggests that CO adsorption on NiY israther strong. Since the isotherms were not reversible no heats could be evaluatedbut a nieasure of the adsorption strength was obtained by adsorbing known amountsof gas, pumping the system for a measured time, and then redetermining the adsorp-tion isotherm. When 0.6 cm3 g-l was adsorbed on NiY-57 only 0.12 cm3 8-l couldbe removed by pumping for 10 niin at room temperature.When 1.3 cm3 g-' wasadsorbed only 0.4 cm3 g-' was removed by pumping for 45 min at room temperature.All the adsorbed CO could be removed from NiY-20 by pumping overnight at roomtemperature.Fig. 5, in which the amount of CO adsorbed at 0°C and 100 N m-2 is plotted as afunction of the extent of nickel exchange, shows that the nickel ions do not have anabsolute preference for hidden sites, However, at approximately 50 % exchangethe amount of adsorption increases rapidly and this indicates that the first nickelions to be exchanged do have a marked preference for a group of 16 sites that are notaccessible to CO. Most probably these are type I sites32 co ADSORPTION ON ZEOLITE YIf one CO molecule could be adsorbed on every exposed nickel ion, and if everynickel ion in excess of the first 16 were exposed, 3.2 cm3 g-1 would be specificallyadsorbed on NiY-66.The measured specific adsorption was only 2.2 cm3 g-l.This implies that even when all the type I sites are filled not all the remaining nickelions are exposed. At high degrees of exchange some nickel ions must occupy typeI' and 11' sites.CUYCuY-87 rapidly adsorbed the first doses of CO, but a slow adsorption becameapparent at a coverage of - 3 cm3 g-'. Equilibrium then took up to 3 days to beattained. A second series of measurements was made in which each adsorption wasterminated after 30min, and a fresh dose of gas was then admitted. In this way" non-equilibrium isotherms " were determined for CuY-87 outgassed at 330°C,440°C and 640"C, and also for CuY-32 and CuY-50 outgassed at 350°C.Theisotherms are shown in fig. 9.FIG. 9.-Adsorption of CO on CuY: 0, "30min " 0°C isotherm on CuY-32 ; 9, " 30 min "0°C isotherm on CuY-50; 0, " 30 min" 0°Cisotherm on CuY-87 outgassed at 350°C; 6," 30 min " 0°C isotherm on CuY-87 outgassed at440°C ; Q, " 30 min " 0°C isotherm on CuY-87outgassed at 640°C ; (3, equilibrium measure-ments on CuY-87 outgassed at 350°C. Thedashed line is the 0°C isotherm onThe curved isotherm and increased adsorption, relative to Nay, show that inCuY-32 cupric ions are accessible to CO molecules. The specific adsorption is 15 %of that expected for a random distribution of cupric ions, or 6 % of that expected ifall the cupric ions were accessible. Thus the behaviour of Cu2+ is intermediatebetween that of Mn2+ and Zn2+. However, fig.5 shows that the change in adsorptionproperties at 50-60 % exchange is not as marked as for MnY. Therefore there is nostrong evidence that the inaccessible cupric ions in CuY-32 have a strong preferencefor type I sites. At higher degrees of exchange more Cu2+ ions are exposed to CO.The equilibrium isotherm for CuY-87 (outgassed at 350°C) implies that at equilibriumat least 10.4 Cu2+ ions per unit cell are accessible to CO. If all ions in excess of thefirst 16 were exposed, 8.4 Cu2+ ions would be accessible to CO. A more detailedinterpretation of the results in terms of site occupancy is complicated by the slowadsorption.C.CO ADSORPTION ON CeYCO adsorption was measured on CeY-40, CeY-71 and CeY-80. Representativeisotherms are shown in fig. 10 and 11. Adsorption on CeY-40 and CeY-71 waT. A . EGERTON AND F . S. STONE 33reversible and rapid. Heats of adsorption were determined for CeY-40 and CeY-71evacuated at 350°C and for CeY-40 evacuated at 460°C (table 5). No further changein the adsorption properties of CeY-40 occurred when the evacuation temperature wasincreased to 615°C. For both CeY-40 and CeY-71 the adsorption heat falls rapidlywith increasing coverage, and at pressures greater than 250 N r r 2 the adsorptionisotherm falls below the corresponding NaY plot. Both these facts indicate thatvery few cerium ions interact with adsorbed CO.The lowered adsorption relativeto NaY probably results from a depletion of the accessible sodium ions.0.40.3rlI M"EFIG. 10.-Adsorption of CO at 0°C on Cc '-40 2and CeY-71: 8, CeY-71 outgassed at 350°C ; 2 c.3(3, CeY-40 outgassed at 350°C; 0, CeY-40 out-gassed at 460°C ; 0, CeY-40 outgassed at 615°C. 4$3 p 0.1' 3 c.5 1.0pressure/kN m-2pressure/kN m-z1-25FIG. 11.-Adsorption of CO at 0°C on CeY-80:0, outgassed at 350°C for 16 h ; (>, outgassed at350°C for a further 32 h ; 0, outgassed at 460°Cfor 18 h ; 0, outgassed at 740°C for 14 h34 co ADSORPTXON ON ZEOLITE YCeY-80 outgassed for 16 h at 350°C adsorbed 3.5 cm3 g-l at 400 N in-’ and 0°C ;evacuation for a further 32 h led to an adsorption of 4.2 cm3 g-l. Neither isothermwas thermodynamically reversible but subsequent outgassing at 460°C led to areversible isotherm.The adsorption at 4-00 N m-’ was then 3.9 cm3 g-l. Thisdecreased when the outgassing temperature was raised to 740°C. The specificadsorption on CeY-80 is unexpected since Rabo and Lo-workers l5 found no cation-specific band in the spectrum of CO adsorbed on Lay. However, recent diffractionstudies have shown that there are differences between the cation distribution indehydrated cerium and lanthanum zeolites.’ Also the CeY-80, unlike the CeY-40and CeY-71, was prepared at 100°C and Sherry l6 has found that lanthanum exchangeat high temperatures is not fully reversible. The unusual adsorption behaviour ofCeY-80 may be a consequence of irreversibly exchanged ions.TABLE 5.-HEATS OF ADSORPTION OF co ON CeYheat of adsorption/kJ mol-1CeY-40 CeY-40 CeY-7 1outgassed outgassed outgassedcoverage/crnJ g-1 at 350°C at 460°C at 350°C0.01 37+ 1.0 42+ 2.0 40+ 2.00.05 29+ 1.0 27.5L4 2.0 42.5f 3.00.10 273- 1.0 28+ 2.0 29.5 *0.20 33k2.00.30 29f 2.00.40 28.5k2.0* This value is based on measurements made at only two temperatures.DISCUSSION1. SITE PREFERENCES OF DIFFERENT CATIONSAll divalent-ion-exchanged fornis of Zeolite Y examined show evidence of specificadsorption of CO. The extent of specific adsorption is variable and reflects thedistribution of cations between the different sites.We have indicated previously Othat these sites fall into three categories. Type I sites are inaccessible to all adsorbedmolecules.Cations in a second group are only accessible to molecules such as water(but not CO) which can enter the sodalite units. Cations in sites I1 and IIo interactwith all molecules which can enter the zeolite cages.Adsorption on low-exchanged zeolite is diagnostic for the distribution of multi-valent cations between accessible (supercage) and inaccessible (type I and sodalite)sites. An earlier studyio showed that Ca2+ ions at exchange levels below 50 %have an absolute preference for hidden sites. The present work shows that for otherdivalent ions the preference, although not absolute, is also strong and increases in theorderZn2+ < Cu2+ < Ni2+, Mn2+, UOZf < Ce3+, Ca2+.A more subtle problem concerns the distribution of the hidden cations betweenhexagonal prisms and the sodalite units.The most probable cause of a markedchange in adsorption properties at - 55 % exchange is preferential occupation of the16 hexagonal prisms per unit cell. Preferential occupation of each of the 8 sodaliteunits by only 2 divalent ions is unlikely since it is known that each sodalite unit canaccommodate at least 3 Ce3+ ionsa2 A compromise explanation that there are(16-x) divalent ions in the hexagonal prisms and x divalent ions in the sodalite unitsis inconsistent with the results of Dempsey and Olson.” These authors found thatin divalent X and Y zeolites there are two type I’ sites occupied for every empty site IT. A. EGERTON AND F. S . STONE 35We therefore conclude that the marked change in the adsorption characteristics ofMnY and NiY zeolites at 50-60 % exchange indicates that the cations mainly occupytype I positions at low degrees of exchange. The slow adsorption on CuY makesinterpretation of site preference difficult, but the adsorption results do not stronglysuggest that type I sites are occupied.Ions in the sodalite units are stabilized by small amounts of residual water, andraising the outgassing temperature removes this water.For ZnY and BaY theadsorption is more sensitive to the outgassing temperature than is the adsorption onNiY-46 and Cay-43. This indicates that for hidden zinc and barium ions thedistribution between the sodalite units and hexagonal prisms is more in favour of thesodalite sites than is the case for calcium and nickel ions.At 57 % exchange there are sufficient divalent cations to fill all 16 hexagonalprisms.If all type I sites are filled any further cations must occupy either sodalitesites or accessible sites. The adsorption results indicate that beyond 57 % exchangethe tendency to occupy the accessible sites decreases in the orderCa2+ > Ni2+ > Co2+.Of the ions investigated, only Ba2+ and U0$+ are sufficiently large to suggest thatthey might be excluded from the inaccessible sites for steric reasons. Barrer, Reesand Davies l8 found that about 16 sodium ions per unit cell of NaY cannot bereplaced by barium ions (Y = 0.134 nm). This suggests that at 25°C Ba2+ ionscannot enter the internal sites. However, since all the sodium can be replaced bypotassium (Y = 0.133 nm) we cannot conclude for certain that steric factors are thedominant ones, even for these three large ions.Since all the transition metal ionsare much smaller than barium it is clear that factors other than ionic radius areimportant in determining their site preferences.These factors include the following : (a) the difficulty of local charge compensationfor multivalent ions; (6) the need to attain a suitable coordination (in some casesby the addition of ligands from water); (c) the need to minimize the electrostaticenergy of the system (direct cation-cation repulsion must be avoided) ; (d) covalency,and directed bonds ; (e) crystal or ligand field stabilization.It would be naive to expect any single one of these causes to determine universallythe cation site preferences in Zeolite Y.Their relative importance will vary with thenature of the ion. Factors (a), (6) and (c) will all be influenced by the presence ofhydroxyl groups formed by the reactionM2+ + H20 -+ MOH+ + H+and the extent of this hydrolysis reaction will depend on the polarizing power of M2+.Since the formation of an MOH+ group stabilizes ions in the sodalite sites, the occupa-tion of sites I' and 11' will also depend both on the polarizing power of the charge-balancing cation and on the outgassing temperature. Crystal field stabilizationenergy will contribute to the stability of tetrahedral Co2+ in the sodalite sites, butcannot explain the affnity of Zn2+ ions for the sites.The distribution of the zincions must be a consequence of the tendency of Zn(I1) to form tetrahedrally directedcovalent bonds, a tendency which is reflected in the structure of many Zn(I1) com-pounds.2. COMPARISON OF THE SITE PREFERENCES WITH THOSE DEDUCED USINGOTHER TECHNIQUESIn this section the site preferences deduced from adsorption experiments are com-For the ions discussed in this paper X-ray diffraction results exist only for nickel,pared with the results of spectroscopic and X-ray diffraction experiments36 co ADSORPTION ON ZEOLITE Ycerium and copper. The nickel ions in a single crystal sample of nickel faujasite[Ni2,Ca,(A10z),,(SiOz),,,] occupy two-thirds of the type I sites and the remainingnickel ions are distributed between sites I1,II’ and I’., In NiY [Nil,Naz,H,(AI02),,(Si02)136] outgassed at 300”C, 10 nickel ions were found to occupy site I, and whenthe evacuation temperature was raised to 600°C the number of nickel ions in site Irose to 12.18 Both sets of results are in fair accord with the present conclusionsbased on CO adsorption.In dehydrated cerium faujasite the cerium ions occupy the I‘ positions.lg InCeX heated in Nz the ions mainly occupy site 1’, but site I and perhaps a site in thecentre of the ring of 12 oxygen atoms which joins adjacent supercages are alsooccupied.2 The adsorption results are consistent with the cerium ions in CeY-40and CeY-71 occupying site 1’.The results for CeY-80 suggest that a few cerium ionsare accessible to adsorbed CO.These accessible cerium ions, perhaps in the site be-tween the supercages, may correspond to the ions which Sherry found to be irreversiblyexchanged when LaY was prepared at high temperatures.The diffraction results for dehydrated Cu, 6Na24Y zeolite 2o indicate that copperions occupy both I and I’ sites. E.p.r. studies 21e 22 indicate that cupric ions occupytwo different environments in Zeolite Y, but no author has yet unambiguouslydefined these environments in terms of the known sites in the zeolite framework.The present CO measurements indicate that type I sites are not exclusively occupiedin either CuY-32 or CuY-50 and are consistent both with the X-ray diffraction andthe e.p.r. studies. The results also show that contrary to the assumption ofRichardson 23 not all the Cu2+ ions are in the supercages.Barry and Lay have measured e.p.r.spectra of Mn2+ probe ions in Zeolite Y.They deduced that in MnY-1 dehydrated at 300°C or above the Mn2+ ions are insites 11, I’ or 11’. Our results indicate that in MnY-32 there are occasional ions insite 11, but mostly the Mn ions are in internal sites, probably site I. The comparisonis not altogether valid because of the widely different concentrations of Mn2+ ionsin the two cases. In the e.p.r. experiments with dilute MnY-1, for example, the Mnions might be associated with a very few 6-membered rings rich in aluminium. Also,the curved heat plot for CO on MnY-75 indicates that the number of accessible Mnions may be temperature-dependent. As already pointed 0utY3 the various factorswhich determine site preference are very finely balanced in this case.In ZnY-67,Barry and Lay deduced from Mn2+ probe spectra that Zn2+ ions occupy sites 11,11‘ and 1’, and that the relative preference of the two ions for site I as compared withsite I1 or similar sites was Mn2+ > Zn2+. Our CO data are in full agreement withthis conclusion, although we doubt whether Zn2+ ions totaZZy shun site I.3 Amongstthe divalent ions we have investigated at - 30 % exchange, Zn2+ has the greatesttendency to show specific adsorption of CO and Ca2+ the least (fig. 2). At highdegrees of exchange, however, the amounts of CO adsorbed on ZnY and CaY aresimilar. For instance, the specific adsorption of 2.73 cm3 8-l on ZnY-62 at 0°Ccompares with 2.5 cm3 g-l on Cay-62, as interpolated from our data on Cay-54and CaY-64.1° This further emphasizes the point that comparison of behaviour atone degree of exchange is not necessarily a good index of relative site preference atsome other, widely different degree of exchange.measured the i.r. spectrum of CO adsorbed on Bay-80and concluded that the first 16 barium ions did not have an absolute preference forsite I, perhaps because of steric effects. However, the adsorption was less thanwould occur if all the Ba2+ were accessible to adsorbed molecules (16.4 instead of22.4 sites per u.c.).The present results confirm that barium is not totally excludedfrom inaccessible sites.Rabo and co-workerT. A . EGERTON AND F . S. STONE 37Rabo et aL9 found that the intensity of the infra-red band for CO on COY wasmuch less than that which would be expected if all Co2+ ions in excess of the first 16were to occupy sites accessible to CO.This result is confirmed in the present workand is attributed to cobalt having a high preference for the I' and 11' sites within thesodalite cages. Ions in these sites would be in quasi-tetrahedral coordination. Sucha coordination is indicated by reflectance spectra and by bulk magnetic suscepti-bility measurements * on dehydrated COY.3. THE NATURE OF THE co ADSORPTIONWe have previously shown that the presence of accessible calcium or zinc ionscauses a marked increase in the heat of adsorption of CO on Y Thepresent results for MnY and BaY confirm that divalent cations in general cause thisincreased adsorption heat.In fig. 12 the low coverage heats of CO adsorption onFIG. 12.-The relationship between the isostericheat of CO adsorption and (1) the i.r. stretchingfrequency of adsorbed CO (open circles) (2) theelectrostatic field of the charge-balancing cation(filled circles)u/cm-2180 2200 2220I I 12 0 I 1 I2 .o 3.0 4 .O10-lo fieldlV m-1Nay, Cay, ZnY and BaY are plotted as a function of the electrostatic field experiencedby the CO molecule. The field was calculated without taking shielding effects intoaccount and the distance of CO from the cation was taken as equal to the ionic radiusof the cation plus the van der Waals radius of CO (0.15 nm). Although the model iscrude, it is evident that the heat of adsorption is proportional to the field strengthexperienced by the CO molecule. All the points except that for CeY lie close to astraight line.It may be that the effective charge on the cerous ion is reduced by aprocess of the kindH20 + O:;mework + Ce3+ + Ce(OH)2+ + OH-.The fit in fig. 12 for CeY would be much more satisfactory if the charge is taken as+2 instead of +3. Since CO has negligible permanent dipole the increase of AHwith field, F, cannot be ascribed to a classical field dipole interaction pF, nor wouldthe classical polarization contribution, +aF2, to the energy of interaction lead to alinear increase in AH with F.Angel1 and Schaffer found that when CO was adsorbed on divalent-ion-ex-changed zeolites the CO stretching band occurred at higher than the gas phasefrequency.The shift, Av? was proportional to the electrostatic field acting on th38 co ADSORPTION ON ZEOLITE YCO niolecule. They suggested that under the influence of the eIectrostatic fieldthere was a transfer of the electrons in the lone pair orbit of the C atom towards thecation of the zeolite ; this process is analogous to the first stage in the formation of a~ - - n bond in a metal carbonyl. The good correlation between Av and AH (fig. 12)suggests that this mechanism also contributes to the high heat of adsorption onzeolites.Angell and Schaffer4 found an anomalous value for Av of CO on NiY andattributed it to the influence of the partially filled d orbitals of Ni2+. The anomalousstrength of the adsorption of CO on NiY which we observed may also be due to thesed orbitals, perhaps because they allow a small degree of back-bonding to occur.On both CuY and NiY samples a slow adsorption occurred.This probablycorresponds to an increase in the number of adsorption sites and is caused by a migra-tion of cations from inaccessible sites in the structure. Gallezot, Ben Taarit andImelik l8 report that NO and NH3 adsorption lowers the number of nickel ions insite I of NiY. Carbon monoxide, unlike ammonia and nitric oxide, is unable toenter the sodalite units of the zeolite structure.lO Therefore the migration of nickelions from inaccessible sites must be induced indirectly; perhaps the CO adsorptionperturbs the environment of the nickel ions by causing a distortion of the zeoliteframework. In NiY most of the inaccessible ions are in the hexagonal prisms, whilstfor CuY most of the inaccessible ions are in sodalite units. Gallezot, Ben Taaritand Imelik 2o have shown that for CuY molecules such as pyridine and butenereadily induce migration of Cu2+ ions from site I' into the supercages. Hence theinaccessible ions in NiY are more completely hidden and this explains why the slowmigration occurs to a much smaller extent in NiY than in CuY.We thank Dr. I. M. Rouse for many helpful discussions.D. H. Olson, J. Phys. Chem., 1969, 72,4366.F. D. Hunter and J. Scherzer, J. Catalysis, 1971,20,246.T. I. Barry and L. A. Lay, J. Phys. Chem. Solids, 1968,29,1395.C . L. Angell and P. C. Schaffer, J. Phys. Chem., 1966,70,1413.K. Klier and M. Ralek, J. Phys. Chem. Solids, 1968, 29, 95.T. A. Egerton, I. M. Rouse and F. S. Stone, to be published.W. N. Delgass, R. L. Garten and M. Boudart, J. Phys. Chem., 1969,73,2970,J. A. Rabo, C. L. Angell, P. H. Kasai and V. Schomaker, Disc. Faraday SOC., 1966, 41, 328.* T. A. Egerton, A. Hagan, F. S. Stone and J. C. Vickerman, J.C.S. Faraday I, 1972, 68,723.lo T. A. Egerton and F. S. Stone, Trans. Faraday Soc., 1970,66,2364.l1 E. Dempsey and D. H. Olson, J. Phys. Chem., 1970,74,305.l2 L. H. Ahrens, Geochim. Cosmochim. Acta, 1952,2, 155.l3 D. J. C. Yates, Canad. J. Chem., 1968,46, 1695.l4 R. M. Barrer and R. M. Gibbons, Trans. Faraday Soc., 1963,59,2569.l5 J . A. Rabo, C. L. Angell and V. Schomaker, Proc. 4th Int. Congress Catalysis, Moscow, 1968l6 H. S. Sherry, J. Colloid Interface Sci., 1968, 28, 288.l 7 R. M. Barrer, J. A. Davies and L. V. C. Rees, J. Inorg. Nuclear Chem., 1968,30, 3333.P. Gallezot, Y. Ben Taarit and B. Imelik, J. Catalysis, 1972, 26, 481l9 D. H. Olson, G. T. Kokotailo and J. F. Charnell, J. Colloid Interface Sci., 1968, 28, 305.2o P. Gallezot, Y. Ben Taarit and B. Imelik, J. Catalysis, 1972, 26, 295.21 J. Turkevich, Y. Ono and J. Soria, J. Catalysis, 1972, 25, 44.22 H. B. Slot and J. L. Verbeek, J. Catalysis, 1968, 12, 2116.23 J. T. Richardson, J. Catalysis, 1967, 9, 178.24 T. A. Egerton and F. S. Stone, J. Colloid Interface Sci., 1972, 38, 195.(Akademiai Kiado, Budapest, 1971), Vol. 11, p. 96 (Paper 54)

ISSN:0300-9599    

DOI:10.1039/F19736900022

出版商:RSC    

年代:1973

数据来源: RSC

摘要:

Magnetic Studies of ZeolitesPart 2.-Magnetic Properties of NiA, Nix and NiYBY T. A. EGERTON* AND J. C. VICKERMAN-~School of Chemistry, The University, Bristol BS8 1TSReceived 10th July, 1972The bulk susceptibilities of nickel-exchanged forms of zeolites A, X and Y measured over a widerange of temperature show that in the fully hydrated zeolites the nickel ions are octahedrally co-ordinated. On dehydration, the octahedral coordination is retained by the nickel ions of NiA andNiCaX. However, highly dehydrated NiA was found to be unstable.For Nix and NiY, the reciprocal susceptibility passed through a maximum when the number ofresidual water molecules were 4 to 6 per Ni2+ ion. This is believed to arise from a distorted octa-hedral or pentahedral coordination around the Ni2+ ions, where some of the ligands are residua1water molecules and the other ligands are lattice oxygen ions.Further dehydration causes a decrease!in the reciprocal susceptibility, this being attributed to many of the Ni2+ ions being in tetrahedralcoordination, where three of the ligands are lattice oxygen and the fourth is derived from residualwater. A tendency for Ni2+ to occupy the hexagonal prisms in preference to sites I’ and II’ isevident in dehydrated CaNiX but not in Nix or NiY. The Ni2+ ions may move out of the hexagonalsites of dehydrated Nix, NiY and CaNiX at T> 400 K ; this behaviour is reversible.We suggest that where most of the Ni2+ ions in a dehydrated zeolite are tetrahedrally coordinated,the nickel ions are less susceptible to reduction.Extensive studies of the catalytic properties of synthetic zeolites have prompted anincreased interest in their solid state properties. Catalytic studies of the nickelexchanged forms I have been accompanied by structural studies using X-ray diffrac-tion,2 electron paramagnetic re~onance,~ and reflectance spectro~copy.~ In Part 1,the utility of magnetic measurements for the study of Co2+ site preferences in zeolitesA and Y was dem~nstrated.~ The magnetic properties of Ni2+ ions are also sensitiveto environment, so bulk susceptibility measurements may be used to study the sitepreferences of nickel ions.This paper describes studies of the bulk magnetic sucept-ibilities of four nickel exchanged zeolites, NiY-66, Nix-52, NiCaX-21 and NiA-62$.Recent studies have suggested that the distribution of silicon and aluminium ionswithin the aluminosilicate framework influences the location of the charge-balancingcations.6 By comparing the magnetic properties of zeolites A, X and Y the influenceof these three frameworks on cation coordination has been examined.The influenceof other cations upon the location of nickel ions was investigated by comparing theproperties of nickel exchanged forms of NaX and CaX, and for all four zeolites the* present address : Department of Chemistry, Makerere University, P.O. Box 16020, Kampala,Uganda. t present address : Department of Chemistry, University of Manchester Institute of Science andTechnology, Manchester M60 1QD.1 The nomenclature used throughout this paper may be described by the following examples :NiY-66 describes a Y zeolite in which 66 % of the charge-balancing sodium ions have been replacedby nickel ions.The nomenclature NiCaX-21 is used to designate an X zeolite in which all the sodiumions have been replaced by calcium ions, 10.5 % of which have subsequently been replaced by nickelions (i.e., the number of nickel ions is 21 % of the initial sodium ion concentration).340 MAGNETIC STUDIES OF ZEOLITESinfluence of water on cation location was investigated by making measurements atdifferent degrees of dehydration.3 0 -v) 25-8*- ii 20- s 1 5 -10-5 -EXPERIMENTALThe ion exchanged zeolites were prepared from NaA (BDH package No. 5816.2/99765),NaX (BDH No.090530) and NaY (Union Carbide SK-40 Lot No. 51-31). CaX wasprepared by percolating calcium chloride solution through NaX at room temperature. Toensure that nickel entered the zeolite solely by ion exchange, and not as a result of nickelhydroxide precipitation, the parent materials were washed with a sodium acetatelacetic acidbuffer (pHw 5.5) before use in an ion-exchange reaction. A small amount of this buffer wasalso added to the first batch of the nickel sulphate solution used for the exchange of thesodium zeolites. (The NiCaX was prepared without the use of buffer.) Ion-exchangedsamples were washed, dried at 343 K, and stored over an aqueous solution of saturatedammonium chIoride at 298 K.The exchanged zeolites were analyzed for nickel (gravimetrically as the dimethylglyoximecomplex) and for sodium (by flame photometry).For the X and Y zeolites, the closecorrespondence between the decrease in sodium content and increase in nickel contentindicated that a simple ion-exchange process had occurred. X-ray diffractometer and lowtemperature adsorption studies indicated that for these samples no loss of structure resultedfrom the exchange process. The analysis of NiA was consistent with the formula :3.73 Ni2+, 3.42 Na+, 1.12 (H30+) [(Si02)12(A102)l,J12-, 38 H20.Oxygen adsorption at 195 K, measured after treatments corresponding to successive dehydra-tion stages used in the magnetic studies, indicated that NiA was not stable to dehydration.The magnetic susceptibilities of the nickel zeolites were measured, as a function oftemperature, using an enclosed Gouy Balance.Measurements were made on the fullyhydrated samples and at several different degrees of dehydration. Controlled dehydrationwas carried out as b e f ~ r e , ~ and in fig. 1 the extent of dehydration is plotted as a function of351, I I I Iocl 160 2 0 0 3GO 400T/"CRG. 1.-Dehydration as a function of temperature. 0, Nix-52 ; El, NiCaX-21; A, NiY-66;+, NiA-62. The open symbols represent dehydration in a helium atmosphere ; the filled symbolsrepresent dehydration in uacuo.the dehydration temperature for all four samples.for the diamagnetism of the aluminosilicate framework, and for the effects of a small amountof ferromagnetic impurity .To study the effects of hydrogen upon the magnetic properties of Nix and NiCaX, thezeolites were contacted with - 30 Torr (5 kN m-2) of hydrogen. In successive experiments,NiCaX was exposed to hydrogen at 523 K, 573 K, 623 K and 693 K.After each exposurethe hydrogen was pumped away before the susceptibility curve was measured. Similarly,Nix was exposed to hydrogen first at 573 K, then 643 K and then 703 K.Measured susceptibilities were correcteT. A. EGERTON A N D J . C. VICKERMAN 41n Y >RESULTSFig. 2 and 3 show representative plots of susceptibility as a function of tempera-ture. At each stage of dehydration the temperature dependence of the nickel ionsusceptibilities could be represented below 400 K byx = c / ( T + ~ ) (1)where C is the Curie constant, related to peff by 8C = p&, and 6 is the Weiss constantwhich reflects the degree of magnetic exchange interactions within the system. Valuesof peff and 6 derived from the susceptibility plots are listed in table 1.1/ I// ‘8// / mT Kweight loss.FIG.2.-Reciprocal susceptibility of Ni2+ in NiA and Nix. 0, Fully hydrated Nix ; 0, NiA, 5.2 %/’ 4 2-c / *FIG..4 ro M6 100\ F43.-Reciprocal loss (final dehydration).For NiX(IV), (V) and (VI), NiCaXO and (VI), and NiY(VI) the susceptibilitycurves changed their temperature dependence above 4-00 K. The changes, whichwere reproducible, were in a direction consistent with higher magnetic moments, andin some cases the susceptibilities returned to their original curves at about 600 K.The derived values of peff, together with the temperature ranges to which they refer,are listed in table 142 MAGNETIC STUDIES OF ZEOLITESThe susceptibility curves for NiA obeyed the Curie-Weiss Law only for the hyd-rated zeolite and for the first four dehydrations.Further dehydrations led to sigmoidsusceptibility curves. Since low temperature adsorption measurements, measuredafter high temperature outgassing, indicated breakdown of the zeolite, the sigmoidcurves may be a consequence of the destruction of the framework. Therefore, onlythe results up to NiA(1V) are presented.TABLE 1 .-MAGNETIC PROPERTIES OF NICKEL-EXCHANGED ZEOLITES AT DIFFERENT DEGREES OFDEHYDRATIONzeoliteNiA(H)NiA(1)NiA(I1)NiA(II1)NiA(1V)NiY(H)NiY(1)NiY(I1)NiY(II1)NiY(1V)NiY(V)NiY(V1)Nix(€€)NiX(II1)NiX(IV)NiX(V)NiX(V1)NiCaX(H)NiCaX(I)NiCaX(I1)NiCaX(II1)NiCaX(IV)NiCa(v)NiCaX(V1)N(HZO/Niz+)11.19.17.14.22.615.212.89.35.22.10.650.013.74.92.70.750.032.628.220.49.15.61.60.0Pert3.293.173.313.263.243.163.153.232.983.263.323.483.162.923.313.373.513.263.203.223.133.223.243.14e- 120-500remarks16163502226520162 T>400 K, ~ ~ f f - 4 .25 T>400 K, ,uu,ff-3.7 ; decreases at 600 K5 T>400 K, 3.9 ; decreases at 600 K0-9-9- 10-9- 10 T>350 K, p , ~ p 4 . 0 ; decreases at 500 K- 10 T>350 K, bff-3.9It is important to compare the susceptibilities of the four nickel zeolites as afunction of dehydration.However, for each sample the measured susceptibilities areinfluenced to different extents by exchange interactions between the nickel ions.(This is reflected in the different 8 values.) For a valid comparison, the effects ofexchange interactions must be eliminated using the equation :where Xcorr is the susceptibility corrected for the effects of exchange interactions, x isthe experimental susceptibility of the nickel ions (after allowance has been made forthe diamagnetism of the alumino-silicate framework) and y = e/C.Elimination of effects of exchange interactions also facilitates comparison ofexperimental and theoretical susceptibilities. The theoretical values for nickel ionsin cubic octahedral and tetrahedral sites may be calculated using equations given byLotgering.' Both theoretical and experimental values are presented in fig. 4a-4dT .A . EGERTON AND J . C . VICKERMAN 43where the experimental values are plotted as a function of N, the number of watermolecules per nickel ion. The calculation of N assumes that the final dehydrationof each series removed all the remaining zeolitic water. If small amounts of waterare retained (probably less than 1 water molecule for every two sodalite cages), thevalues of N corresponding to the final stages of dehydration will be slightly low.The plots for NiA and NiCaX are similar and there is little variation of x;kr withL- 0 1 0 xO 2 b 16 ” I; A I ; ’ 0N(Hz 0 /Ni2 +) N(H,O/Ni”+)(4 (4FIG. 4.-Corrected reciprocal susceptibility isotherms for a, NiA-62 ; b, NiY-66 ; c, Nix-52 ; d,NiCaX-21 at : 0 , 100 K ; 0, 200 K ; 0, 300 K ; 8, 400 K.The calculated susceptibilities fortetrahedral Ni2+ are plotted at the left and those for octahedral at the right of each figure44 MAGNETIC STUDIES OF ZEOLITESchange of N ; the susceptibility values cluster around the theoretical value for octa-hedral nickel. For high values of N , susceptibilities of NIX and NiY also clusteraround the octahedral value, but for 6>N> 5 the values pass through a markedmaximum. For N<3, &f, falls sharply, but it never drops as far as the theoreticalvalue for tetrahedral nickel.The variation of magnetic moment with dehydration shows a similar trend.Forall four samples, the initial moments of 3.1-3.2 pB are close to values expected foroctahedral nickel. At N = 5, the moments of Nix and NiY fall to 2.9 pB but thereis little change in the moments of NiA and NiCaX. As dehydration proceeds stillfurther, the moments of Nix and NiY, but not of NiCaX, rise to 3.4-3.5 pB, which isstill lower than the value of about 4.0 pB expected for cubic tetrahedral nickel. How-ever, when N< 2 or 3, values of peff in excess of 3.7 pB were obtained for Nix, NiY andNiCaX in the temperature range between 400 and 600 K.Exposure of NiCaX to hydrogen at T<623 K caused a gradual decrease in x - lwith each treatment (table 2). The magnetic moment rose slightly, to 3.4 pB at theTABLE 2.-RECIPROCAL SUSCEPTIBILITY OF Ni2+ AS A FUNCTION OF TEMPERATURE AND THEEFFECT OF EXPOSURE TO HYDROGEN ON THE MAGNETIC PROPERTIES OF Nix-52 AND NiCaX-21x-1 31-1 2-1 x-lNiz+ environment details of treatment /tppr 100 200 300 400 remarkstetrahedral theoretical valueoctahedral theoretical valueNiX-52(VI) after final dehydrationNiX-52(VI) + H z30 Torr for 3 h at 583 K30 Torr for 3 h at 643 K10 Torr for 3 h at 703 K15 Torr for 5 h at 523 K50 Torr for 5 h at 573 K50 Torr for 5 h at 623 K30iTorr for 15 h at 683 KNiCaX-21WI) after final dehydration+ H267 85 113 14577 154 231 2933.51 70 133 198 258 T>400K, peff4.2;3.51 68 132 198 258 T>400K, 1.~eff3.9;3.51 66 132 198 258 T>400K, peff3.9;3.14 73 153 231 295 T>350K, ~eff3.9decreases 600 Kdecreases 600 K3.62 65 127 188 2503.39 75 146 216 2873.30 76 153 229 304ferromagnetic Tc = 545 K, see fig.5.3.35 52 140 214 273 T>320K, peff3.9FIG. 5.-A plot of magnetisation against T, for Nix-52 and NiCaX-21 treated in Hz at 703 and683 K respectivelyT . A. EGERTON AND J . C . VICKERMAN 45first treatment and then stayed steady. After H2 adsorption at 693 K, there was alarge decrease in 2-l and a plot o f 0 (magnetization) against Tgave a Curie tempera-ture of 545 K (cf. 630 K for massive nickel), see fig. 5. Similar treatment of NIX atT = 573 K and 703 K did not appear to reduce the nickel ions significantly; x-l at300 K fell from 198 to 188 and peff rose to 3.62 pB.DISCUSSIONThree different types of site are available to the charge-balancing cations ofdehydrated faujasite-like zeolites ; each is associated with a different coordination.To identify these sites we use the nomenclature described previously.' Site I islocated within the hexagonal prisms.When nickel ions occupy this site they slightlydistort the aluminosilicate framework and attain a coordination that is very close tooctahedral.2 Sites I1 and IIo, in the hexagonal windows between sodalite units andsupercages, have trigonal symmetry. Ions within the sodalite units (site I' and 11'),although nominally in trigonal symmetry, attain a near-tetrahedral coordination if,in addition to the oxygen ligands of the aluminosilicate structure, there is a fourthoxygen, near the centre of the unit, to which they can coordinate.This fourthoxygen may be from a residual water molecule, or from a hydroxyl group derivedfrom such a Although similar sites are available in the hydrated zeolites,many cations occupy no fixed positions, but instead '' float " in the zeolitic water.gIn hydrated zeolite A, many cations float in the zeolitic water.g In each unit cellof dehydrated zeolite A there are 8 sites similar to the site I1 locations of zeolites Xand Y, and a further 6 sites near the centre of the rings of eight oxygens which form thewindows between adjacent large cages.l0 In calcium A the first type of site isoccupied preferentially by the six calcium ions.During the early stages of dehydration, when N>6, all the zeolites studied hadmagnetic moments of 3.2 pB and x&tr was close to the expected value for octahedralNi2+.We conclude that the nickel ions are octahedrally coordinated both in thehydrated zeolites and during the initial stages of dehydration. In NiA, the ions canattain octahedral coordination when they form the Ni(H20)g+ complex in the zeoliticwater. In the X and Y zeolites, the type I sites provide a second environment withthe required symmetry. However, Dempsey and Olson conclude from diffractionmeasurements that divalent ions do not occupy site I positions until most of thezeolitic water has been removed.12 Most probably the nickel ions exist as hexaquocomplexes within the large cages of all three zeolites, provided that there are sufficientwater molecules.When Nfell to approximately 5, the magnetic behaviour of NIX and NiY divergedfrom that of NiCaX and NiA.For Nix and NiY an increase in x;&, and an associa-ted fall in peff suggested that the coordination of some or all of the nickel ions hadbeen altered in such a way as to quench the orbital contribution to the magneticmoment. Therefore, the possibility that the nickel ions are distributed between thecoordinations Ni(HzO)2+ and Ni(H20)$+, as was found with ~ o b a l t , ~ is inconsistentwith the magnetic results. Tetrahedral nickel is normally associated with increasedpeff and lowered reciprocal susceptibility, the opposite of the present results. Twopossible explanations exist. First, some five-coordinate species could be formed.Square pyramidal and trigonal bipyramidal Ni(I1) complexes are known, and bothmay exist either as high spin complexes, peff -3.3 pB, or as low spin diamagneticspecies.13* l4 Ni(H20)5-ZOx, (where Ox designates a framework oxygen) is morelikely than Ni(H20)5 because 5-coordinate nickel complexes are usually stabilizedby the presence of at least one rather bulky ligand.A possible species would b46 MAGNETIC STUDIES OF ZEOLITESNi(H20)20~3 in which the three Ox ligands are all from a group of three oxygenatoms associated with one of the windows joining a supercage to a sodalite unit. Asecond possibility is that as N decreases some Ni(H,O), loses its coordinating watermolecules to form Ni(H20),-,0xz. The resulting distortion from pure octahedralcoordination would be expected to lower the orbital contribution to the magneticsusceptibility.It is difficult to distinguish between these two possibilities ; however,it is implicit in both explanations that at least some nickel ions are coordinated byframework oxygens.By contrast, the magnetic properties of NiA show remarkably little change in theregion N = 5, although it is evident that not all the nickel ions can remain in a hexaquocomplex. Nor can the nickel ions retain octahedral coordination by occupying type Isites as there are no hexagonal prisms in NiA. It seems most likely that a mixture ofspecies such as Ni(H,O),, Ni(H20)6-zOxZ and Ni(H,O)Ox,, the relative concentra-tions of the different species depending on the precise degree of hydration at anyinstant. The lowered susceptibility resulting from the presence of Ni(H20) 6-zOxzwould be compensated by a rise in the orbital contribution due to the presence ofNi(H,O)Ox,. Klier and Ralek studied the reflectance spectrum of NiA-29 andfound evidence for a quasi-tetrahedral Ni(H,O)Ox, from N = 2.8 to N = l.L4They suggest that the nickel ions are associated with the six membered windows ofthe sodalite units and project into the large cavities of the zeolite A structure.The magnetic properties of NiCaX also contrast sharply with those of Nix andNiY but, unlike NiA, the nickel ions could retain octahedral coordination byoccupying type I sites.During the final stages of dehydration there is little change inmagnetic properties below 300 K ; this suggests that the nickel ions in NiCaX dooccupy site I.It is significant that the magnetic data suggest that the great majorityof the Ni2+ ions occupy site I positions. It has been suggested, on the basis of electro-static and strain energy considerations, that the preference of divalent cations foroccupation of the site I positions will tend to rise with cation size.12 However, thelarger Ca2+ ions do not appear to block these sites to occupation by Ni2+ ions, indeedcomparison with the NIX data would suggest that the Ca2+ ions block the sites I’ and11’. The e.p.r. results of Barry and Lay l5 offer support to this suggestion. TheCa2+ ions did not prevent Mn2+ ions from entering site I positions in CaX-88 out-gassed at 573 and 673 K.High temperature transitions were observed in the susceptibility data for NiCaX(V)and NiCaX(V1).These data, which lead to increased values for Peff, suggest thatsome of the nickel ions may move out of the hexagonal prisms when the temperatureis raised.For Nix and NiY, dehydration beyond N = 3 results in decreasing x-’ and in-creasing peff. The new reciprocal susceptibilities are well below the expected valuesfor cubic octahedral Ni2+ (see table 2), but are above the values for undistortedtetrahedral nickel. It seems likely that some of the nickel ions attain octahedralcoordination in the hexagonal prisms whilst others are distributed between the sitesof trigonal symmetry (11, ITo) and the sodalite sites in which they can coordinate tothe framework oxygens and a residual water molecules to form an Ni(H20)Ox, species.(Since there are 23 nickels per unit cell of Nix-52 and - 19 per unit cell of NiY-66not all can be accommodated in the hexagonal prisms.) During the later stages ofdehydration, hydrolysis of nickel ions can take place according to the reaction l6Ox2-+Ni2-~+H20-+Ni(OH)++OxH-.Thus the Ni(H,O)Ox, species may be replaced by Ni(OH)Ox, or even by Ni(O)Ox,.The magnetic behaviour of these species will be similar to that of Ni(H,O)Ox,T . A .EGERTON AND J . C. VICKERMAN 47However, unlike the Ni(H,O)Ox,, the Ni(OH)Ox, and Ni(O)Ox, will be relativelyimmobile. The symmetry of the Ni(H20)Ox3 species is trigonally or tetragonallydistorted tetrahedral ; it can be expected to have similar magnetic properties to nickelrhodate in which the nickel ions have similar coordination when the temperatureis below 400 K.17 Above 400 K, the site symmetry of Ni2+ ions in NiRh204 changesfrom tetragonally distorted tetrahedral to cubic tetrahedral and the peff rises from3.4 pB to 4.1 pB.It is possible that a similar transition affects the tetrahedrallycoordinated ions in the zeolite. However, the change of site symmetry in NiRh204results from a rearrangement of four similar ligands (oxygen ions). In the zeolitethere are three ligands of one type (framework oxygens) but the fourth is probably anoxygen from a water molecule or hydroxyl group, and the electrostatic field would notbe homogeneously tetrahedral even if the four ligands had the required geometricalpositions. It therefore seems more probable that in Nix and NiY, as in NiCaX, thechange in magnetic properties is caused by nickel ions moving from the hexagonalprisms to other sites.Since any remaining zeolite water is likely to be coordinatedwith those Ni2+ ions that already occupy sodalite sites there will be little water left tostabilize the nickels that migrate from the hexagonal prisms. Some nickel ions couldoccupy sites I1 and IIo in trigonal coordination to the lattice oxygen (NiOx,) wherethey would be accessible to adsorbed molecules, and be able to take part in catalyticreactions. Such a coordination could also arise in the course of the dehydrationprocedure although it is believed that the number of NiOx, species arising by thisroute would be small since the vacuum conditions used here would be unlikely toremove all traces of zeolitic water.5 If the ligand field around the nickel ions is D31rthe ground state will probably be 3A2.4 In this case, a low magnetic moment andhigh x-l would be expected.The trend of x-l is, however, downward suggestingthat if NiOx, is present the ligand field symmetry is C3". In this case an 3E groundstate is likely, which would make a significant orbital contribution to the suscepti-bility.THE EFFECT OF HYDROGEN TREATMENTHydrogen treatment of Nix-52 led to rather small decreases in reciprocal suscepti-bility. These small changes may be a consequence of the displacement of residualwater molecules from a small number of cations,l' or may be a consequence of thereduction of a very few Ni2+ ions, perhaps located in sites where the framework isparticularly rich in a1uminium.l However, there was no evidence for reduction ofnickel ions to massive nickel crystallites. By contrast, NiCaX became ferromagneticafter hydrogen treatment at 683K.Thus the resistance of the two samples toreduction differs markedly.An examination of the magnetisation curves, at 100 and 310 K, as a function ofmagnetic field strength for hydrogen treated CaNiX (at 623 and 683 K) and NIX (at703 and 643 K) (see fig. 6) highlight the differing reducibility of the two samples.The high magnetisation of CaNiX(H, 683 K) and the small difference between thethe curves at 100 and 310K are characteristic of large nickel cry~tallites.~~ Themagnetisation curves for Nix are inuch lower and there are rather large differencesbetween the magnetisation at 100 and 310 K.Clearly there are a large number ofvery small nickel crystallites. However, although the temperature of H, treatmentis higher for Nix than for CaNiX, it could be argued that the different extents ofreduction are attributable to the differing reduction times. It is thus significant thatthe magnetisation curves demonstrate a clear difference in the reducibility of CaNiXat 623 K and NIX at 643 K. Although the amount of metallic nickel formed i48 MAGNETIC STUDIES OF ZEOLITESCaNiX is small and the particles are small, reduction has clearly occurred, whereasthere is no evidence of reduction of Nix at 643 K.This difference in stability may be a consequence of the different cation distributionin the two zeolites.The susceptibility measurements indicate that in Nix-52 manynickel ions are tetrahedrally coordinated as Ni(OH)Ox,, or related species, in thesodalite cages. Most of the nickel in NiCaX-21 is in the hexagonal prisms, althoughat high temperatures some of the ions may migrate to other sites. Cimino andcoworkers have found that Ni2+ tetrahedrally coordinated at the surface of aluminais much less easily reduced than the corresponding octahedrally coordinated ions. 21It is also pertinent that the Mossbauer spectrum of Fe(1I)Y-65 evacuated at 673 Kshows evidence of ferrous ions in fourfold coordination.22 This zeolite could not bereduced in H2 at 673 K. By contrast, FeCaX, dehydrated at 633 K, did not show thespectrum characteristic of tetrahedral Fe(I1) and this zeolite was reduced in H2 at633 K.23FIGfor0,+ 1.2 .:I I .o * *o * 2 r i , l H ti: 2000 3 0 0 0 4000 5 0 0 0 6000 7CHIG)O'.6.-The variation of magnetisation as a function of magnetic field strength at lo0 and 310KNiX-52 and CaNiX-21 treated in hydrogen (at temperatures shown in brackets). At 100 K :Nix-52 (703 K) ; a, Nix-52 (643 K) ; A, CaNiX-21 (683 K) ; V, CaNiX-21 (623 K). At310 K: 0, Nix-52 (703 K) ; A, CaNiX-21 (683 K); v, CaNiX-21 (623 K).For alumina it was suggested 21 that tetrahedrally coordinated cations were lesssusceptible to reduction due to their tendency to bind their anions more closely thanoctahedrally coordinated cations.Such an explanation may also be valid in the caseof Ni zeolites since it is likely that H2 can penetrate to cations in both the sodaliteunits and the hexagonal prisms.It is interesting to note that Yates was able to reduce Nix, previously evacuatedat 473 K, in a stream of H2 at 673 K.24 Quite large (d-240A) nickel crystalliteswere formed.Romanowski l9 and Rickert 2o have, however, demonstrated that the reductionprocess is very sensitive not only to temperature but also to H2 pressure and time ofexposure. The present work suggests that the reducibility of Ni zeolites may alsobe very sensitive to the identity of other charge balancing cations present, and to theprecise details of dehydration. This may be the cause of the rather irreproduciblebehaviour that has been found in catalytic studies of the nickel zeolitesT. A . EGERTON A N D J . C. VICKERMAN 49We thank Prof. F. S. Stone for his encouragement and help. We are grateful forthe award of an I.C.I. Fellowship (J. C. V), and for an S.R.C. Research Assistantship(T. A. E.).C. G. Pope and C. Kemball, Trans. Faraday SOC., 1969,65,619.D. H. Olson, J. Phys. Chem., 1968,72,4366.J. A. Rabo, C. L. Angell, P. H. Kasai and V. Schomaker, Disc. Faraday Suc., 1966,41, 328.K. Klier and M. Ralek, J. Phys. Chem. Solids, 1968,29,951.T. A. Egerton, A. Hagan, F. S. Stone and J. C. Vickerman, J.C.S. Faraday I, 1972, 68,723.E. Dempsey, J. Phys. Chem., 1969,73, 3660.F. Lotgering, J. Phys. Chem. Solids, 1962,23, 1153.T. A. Egerton and F. S . Stone, Trans. Faraday Soc., 1970,66,2364.L. Broussard and D. P. Schomaker, J. Amer. Chem. Suc., 1960,82,1041.lo J. V. Smith and L. G. Dowell, 2. Krisf., 1968,126, 135.K. Seff and D. P. Shomaker, Acfa Crysf., 1967, 22, 162.l2 E. Dempsey and D. H. Olson, J. Phys. Chem., 1970,74, 305.L. Samni, Pure Appl. Chem., 1968,17,95.l4 L. Sacconi and I. Bertini, J. Amer. Chem. Suc., 1968,90,5443.Is T. I. Barry and L. A. Lay, J. Phys. Chem. Solidr, 1968,29,1395.l6 J. W . Ward, J. Catalysis, 1969, 14, 365.I’ G. Blasse, Phillips Res. Rep. Supp. No. 3, 1964.l 8 C. Kemball and R. McCosh, Proc. Roy. Soc. A, 1971,321, 249.l9 W. Romanowski, 2. anorg. allg. Chem., 1967,351, 180.’O L. Rickert, Ber. Bunsenges. phys. Chem., 1969,73,331.21 M . LoJacono, M. Schiavello and A. Cimino, J. Phys. Chem., 1971,75,1044, 1051.22 W. N. Delgass, R. L. Garten and M. Boudart, J. Phys. Chem., 1969,73,2970.23 J. A. Morice and L. V. C. Rees, Tram. Faraday Suc., 1968,64, 1388.24 D. J. C. Yates, J. Phys. Chem., 1965,69,1676

ISSN:0300-9599    

DOI:10.1039/F19736900039

出版商:RSC    

年代:1973

数据来源: RSC

摘要:

Sintering Studies on Platinum Black CatalystsPart 1.-Effect of Pretreatment and Reaction on Particle SizeBY THOMAS BAIRD, ZOLTAN PALL j- AND SAMUEL 3. THOMSON *Chemistry Dept., The University, Glasgow GI2 8QQReceived 20th July, 1972Counter claims appear in the literature on the effects of the particle size, sintering and pre-treat-ment of a catalyst on its activity. We have investigated these by studying a platinum black catalystand the way in which helium, hydrogen, air, oxygen and thermal cycling affect the particle size asseen in the electron microscope. By making a logical sequence of experiments it has been possibleto show that the sintering of the catalyst was most affected by the presence of hydrogen which causedconsiderable growth in crystallite size.It thus seems possible that apparent effects of reaction andthermal treatment may have, in previous work, been caused by pre-treatment in hydrogen.Platinum is one of the most commonly used catalysts for hydrocarbon trans-formations. There are differing views in the literature on the extent to which theparticle size or dispersion of platinum may affect its catalytic activity or selectivity.It was found that several reactions such as hydrogenation of cyclopropane,l de-hydrogenation of cyclohexaneY2 hydrogenation of ben~ene,~ were independent of theparticle size of platinum. These reactions were called by Boudart " facile " reactions.'On the other hand, the so-called " demanding " reactions require specific surfacestructures, the concentration of which is dependent on dispersion and surface area ofplatinum.Such reactions are the hydrogenolysis of neopentane and skeletalisomerization by various mechanism^.^.Platinum catalysts of different dispersion were prepared in different ways.Gault observed different selectivity of platinum catalysts containing different amountsof platinum 5 * ; he explained this phenomenon by the different dispersion of themetal in the catalysts. Poltorak on the other hand found that by increasing theplatinum content of his Pt/SiO, catalysts prepared in different ways, only the numberof Pt-aggregates increased but the size of individual metal clusters (or crystals) did not.It was mainly the method of preparation which affected the activity and selectivityof the catalysts used by him.Boudart and co-workers prepared platinum samplesof different surface area by sintering them at various temperatures between 425 and900°C and found different selectivities for hydrogenation and hydrogenolysis of neo-pentane. Their treatment, however, was not uniform for every sample, and involvedheat treatment in hydrogen and/or vacuum or eventually a cyclic treatment withoxygen and hydrogen. The sintering effect nevertheless was attributed exclusivelyto the thermal treatment.Wilson and Hall investigated the behaviour of alumina- and zeolite-supportedplatinum catalysts heated in hydrogen, with and without oxygen pretreatment, atdifferent temperatures. They observed a change in the surface area of the metal,t permanent address : Institute of Isotopes of the Hungarian Academy of Sciences, Budapest114 POB 77.5T. BAIRD, z.PALL AND s. J . THOMSON 51according to X-ray and electron microscopic measurements, and they deduced that“ heat treatment did cause growth in crystal size ”.On the other hand the decrease of area of silica-supported platinum was notaccompanied by growth of crystal size.8In experiments which may be related to those which have been described it wasfound ’* * that the stoichiometry of hydrogen-oxygen titration, used in the determina-tion of metal surface areas,g* lo was affected by sintering the catalyst at differenttemperatures. It was the hydrogen chemisorption which followed the real metalsurface. This point has been investigated by Akhtar and Tompkins 11* l 2 whofound that at elevated temperatures the chemisorbed oxygen may migrate into thebulk of the platinum and can thus be removed incompletely and slowly by subsequenthydrogen treatment.Hence, it cannot be surprising that a controlled oxygen treat-ment of a sintered reforming catalyst may lead to redistribution of the platinumparticles. Such a phenomenon may cause the different catalytic activities ofplatinum samples which have, or have not, been in contact with 0xygen.l’These studies suggest that the subtle relationships between thermal treatment ofplatinum in different atmospheres on the one hand and its surface and catalyticactivity on the other cannot be regarded as having been elucidated. It is for thesereasons that we decided to carry out these investigations.Supported catalysts are as a rule stable and active and the relative instability ofunsupported platinum black may thus offer better opportunities for the study of itschanges during pre-treatment and reaction.It is also advantageous in that thepossible effects of the support,14 which may play an important part in catalysis, canbe eliminated.Sintering of Pd-black has been observed by heating it in hydrogen at 60-90”C.15One of the direct observations which stimulated this part of our study was the factthat during a prolonged period of use the surface area of the same platinum blackcatalyst had decreased by a factor of ten, whereas only an approximately twofolddecrease in its activity was observed in the aromatisation of n-hexane.16EXPERIMENTALPlatinum black was produced by reduction of H,PtCI, by formaldehyde in the presenceof concentrated KOH at 20 to 25OC.l’ The catalyst was washed with distilled water forseveral weeks, then filtered and stored in the air.Its surface area (as determined by B.E.T.method using krypton as the adsorbate) was - 8 m2 g-l.Changes in particle sizes were followed by electron microscopy. The electron micro-graphs were taken as follows : specimens suitable for electron microscopy were prepared byallowing a drop of distilled water with the catalyst in suspension to dry down on to carbon-film specimen grids. The suspensions were previously placed in an ultrasonic bath for 5 sto aid dispersion of the solid. Samples prepared in the dry state gave similar results, butthe adherence of the catalysts to the supporting film was poor and the former method waspreferred.The specimens were examined with a Siemens Elmiskop I electron microscopeoperating at 100 kV.The different treatments of the catalyst have been carried out in a flow type reactor. Allthe different gases were passed over the platinum at a flow rate of about 50 ml/min at about1 atm total pressure except For the air treatment which was applied simply by opening thereactor to the atmosphere. Most of the samples were obtained by a uniforni and controlledtreatment of fresh, unused Pt called the standard thermal cycle (s.t.c.). This consisted of aheating period from room temperature to 360°C within 25-30 min.The temperature wasraised within about 10 min to 330-340°C, then the final temperature of 360°C was adjustedduring the rest of the warming up period. The temperature was then kept constant and 3 hafter the start the heating was switched off. The sample was left to cool in helium flow 52 PARTICLE SIZES I N A PLATINUM BLACK CATALYSTthe cooling period was 45-50 min. The catalyst samples were removed from the apparatusafter reaching room temperature.RESULTSThe electron micrograph of the fresh catalyst shows a very finely dispersedplatinum black, fig. 1. On the contrary, the same catalyst kept for 50 h at elevatedtemperature not exceeding 360°C and used for study of ethylene reactions in heliumshowed remarkable sintering which caused the particle size to grow, fig.2. Thesereactions involved injecting ethylene pulses onto the catalyst via the helium carriergas. The products of reaction were ethane and methane. Before reaction thecatalyst was heated in air and subsequently in hydrogen (s.t.c., the exact treatmentis described in the caption of fig. 6); deactivated catalyst was regenerated with airand then hydrogen treatment at reaction temperatures. Regeneration took placeseveral times in the case of the sample shown in fig. 2.X-ray line broadening measurements confirmed that the phenomenon observedwas actually the growth of the average crystallite sizes not just surface aggregationor welding of smaller particles. The average crystallite sizes determined by electronmicroscopic observation and X-ray line broadening were as follows.crystallite sizes/Aas determined byelectron micrography X-ray diffractionfresh Pt SO+ 30 l l S + Sused Pt 500+ 60 500+ 35In order to investigate the causes of this sintering we decided to treat the platinumblack in different ways which corresponded to the different stages through which thecatalyst had passed in the ethylene experiments.The standard thermal cycle wasused in every case : the different treatments to which the catalysts were exposed areshown in the legends to the corresponding electron micrographs (fig. 3-8).Hardly any sintering could be observed in catalysts which had been heated inhelium or in air. Likewise, heating in vacuum produced pictures identical with fig. 3.On the other hand samples which had been in contact with hydrogen showedsintering to a greater or lesser extent.The effect was more marked when platinumblack had been pre-treated in air at elevated temperature, fig. 5 and 6.Surprisingly, the ethylene treatment caused only minor sintering as can be seenin fig. 7. When the catalyst which had been used for ethylene studies had beentreated subsequently by hydrogen, virtually no further sintering was observed, fig. 8.These results make it clear that the sintering process is connected in some waywith the hydrogen treatment. To elucidate this point further, catalyst samples weretreated with hydrogen for 60 min, but under different thermal conditions.It can be seen in fig. 9 that hydrogen at room temperature caused no sintering :hydrogen at room temperature followed by the standard thermal cycle in heliumcaused some sintering, but only to a small, uniform extent, fig.10. When heatingwas carried out in hydrogen in the third case, fig. 1 1 ? a considerable, but not complete,sintering was observed.The treatment of the sample for fig. 6 under the standard thermal cycle representedone of our standard pretreatments of catalysts for the study of hydrocarbon reactions.It is therefore of interest to compare the catalytic activities, in a microcatalyticreactor gas chromatographic system, of treated and untreated catalysts. The resultsare shown in table 1FIG. 1.-Electron micrograph of fresh, un- FIG. 2.-Electron micrographs of the sametreated, finely divided platinum black.Particle catalyst as fig. 1, but maintained at - 360°Csize, about 1008, : magnification x 48 000. for 50 h and used for ethylene reactions inhelium (regenerated several times with air andhydrogen). Particle size, about 500 8, : magni-fication x 48 000.FIG. 3.-Electron micrograph of platinum FIG. 4.-Electron micrograph of platinumblack subjected to helium flow and standard black. Sample heated in air with subsequent 30thermal cycle : magnification x 48 OOO. min in air, then helium : magnification x 48 000.[To face page 5FIG. 5.-Electron micrograph of platinum FIG. 6.-Electron micrograph of platinumblack : sample heated and subsequently held black : sample heated in air and subsequentlyfor 30 min in helium, then 60 min of hydrogen 30 min in air, then flushed with helium, then60min in hydrogen, then helium: magnifi-cation x48 000.flow, then helium : magnification x 48 000.FIG.7.-Electron micrograph of platinum FIG. 8.-Electron micrograph of platinumblack : sample heated in He, then 11 x 0.5 ml black : previous catalyst, fig. 7, kept for 2 daysethylene slugs passed over catalyst (in helium in air, then s.t.c. : heated and kept for 30 mincarrier gas) : magnification x 48 OOO. in He, then 60min H2, then He: Magnifi-cation x 48 000FIG. 9.-Electron micrograph of platinum FIG. 10.-Electron micrograph of platinum :black : sample exposed to He flow for 10 niin, sample exposed to He 10 min, H2 60 min atthen H2 60 min at room temp., He 3 h, room room temp., He and s.t.c.in He : magnificationtemp. : magnification x 48 000. x48 000.FIG. 11 .-Electron micrograph of platinum FIG. 12.-Electron micrograph of platinumblack : sample exposed to He flow for 10 min, black heated to 600°C in air : magnificationH2 10min at room temp., heated to 360°C inH2 over 25 min, maintained in H2 for 25 min,then s.t.c. completed in He : magnificationx48 000.x 48 000T. BAIRD, z. PALL AND s. J . THOMSON 53TABLE 1 .-PRODUCTS FROM ETHYLENE INTERACTIONS WITH TREATED AND UNTREATEDPLATINUM BLACK CATALYSTScatalyst, 0.06 g Pt ; 0.5 ml pulses of ethylene into 60 ml/min He flow; temp. 360°Cproducts/ %retained oxidisedCZH4 species products pulse no. CH4 C2H6treated catalyst3.4 18.8 60.0 17.8 - 12 1.4 6.1 85.6 6.9 -6 1.3 1.5 97.2 I -untreated catalyst1 16.9 33.4 14.5 22.4 12.82 4.2 16.0 70.8 9.06 1.7 1.8 96.5-- -DISCUSSIONThe electron micrographs show clearly and without doubt that the sintering ofour platinum black had been practically completed during the pretreatment and wassomehow connected with the hydrogen treatment.Different mechanisms can beput forward for the sintering. (a) It could be the result of the heat of reactionbetween oxygen and hydrogen on the platinum surface. (b) It could be attributedto the interaction of hydrogen with the catalyst.It is well known that the platinum is covered with a layer of oxygen when it hasbeen in contact with air. The reaction between hydrogen and oxygen can be due tothe removal of this oxygen layer. Before exposure to hydrogen, the catalyst wasalways flushed with helium, thus no atmospheric oxygen was present.If we assumea monolayer of oxygen on the platinum with 1.1 x 1015 surface atom/cm2, the reactionheat would be enough to melt a considerable fraction of the platinum, provided thatthermal conduction were negligible. Since the hydrogen treatment at room tempera-ture and at 360°C resulted in completely different pictures, and the heating ofhydrogen-treated platinum in helium produced a third variant, this assumption maybe rejected, as the only or the main cause of the sintering.Another possible hydrogen + oxygen reaction might have arisen when the hydrogen-treated samples were taken out of the reactor into the open air. Since no pyrophoricphenomena were observed, and what is more, since the catalysts treated with hydrogenunder different conditions, but having the same concluding period of the thermalcycle in helium, showed different pictures, this suggestion, too, can be regarded asfalse.The sintering phenomena must be interpreted in terms of hypothesis (b), that isit must be connected with the hydrogen treatment itself.When platinum surfaces covered with an oxygen monolayer are contacted withhydrogen, the latter will displace the oxygen on the surface.'-1° This process cantake place at room temperature.This reaction obviously does not cause any sintering.If, however, relatively moderate thermal treatment was applied with the hydrogen,it altered the catalyst structure. It can be seen in fig.1 and 6 that this phenomenonis accompanied by the increase of the crystal size and cannot be regarded only as anagglomerating process.The sintering of Pd-black was explained by " meniscus formation " betwee54 PARTICLE SIZES I N A PLATINUM BLACK CATALYSTindividual crystallites. Although the experiments were done in hydrogen, theauthor has not emphasized the importance of its presence.We regard it as obvious that this process required a continuous hydrogen supply.This is confirmed by comparing fig. 5, 1 0 and 11. When there is a hydrogen supplyat higher temperatures, fig. 5 and 1 1, extensive recrystallization occurs, although thiswas not uniform throughout the catalyst : the process took place mainly at the metalsites attacked originally.On the other hand, when the hydrogen layer was formedat room temperature, then heated in helium, this caused only a small, but uniformrecrystallization to the extent which was allowed by the uniform hydrogen layerformed at room temperature.Obviously, the surface hydrogen at higher temperature starts an interaction withthe catalyst which leads to its " bulk " recrystallization. It can be assumed that thismay start by diffusion of hydrogen into the metal.Radioactive tracer experiments carried out with hydrogen gas labelled withtritium clearly show that at 360°C there are at least two types of hydrogen held by theplatinum black (in the absence of gas phase hydrogen). These species can be trans-formed into each other and could be exchanged with gas phase hydrogen.Catalystsamples treated with radioactive hydrogen in ways analogous to those shown in thecaptions of fig. 6, 9, 1 0 and 11 were still radioactive after taking them out into airwhich must have originated from tritium held by surface or sub-surface 1a~ers.l~l2 have shown that at higher temperatures oxygenmay diffuse into the platinum metal. If this was allowed by an oxygen treatment atelevated temperatures (fig. 6), the structure of platinum already loosened to someextent by oxygen, but not to an extent causing visible sintering-see fig. 4, mayrecrystallize more readily under the effect of hydrogen. In this particular case thesintering may be enhanced by the reaction heat between the oxygen, this time in excessof a monolayer, and hydrogen.It must be pointed out that the sintering process may also take place in an oxygenatmosphere, but it requires more severe conditions. Fig.1 2 shows a sample heatedto 600°C in air : severe sintering can be observed.In the case of untreated catalysis it would be expected that the layer of oxygen ontheir surfaces would be susceptible to removal by hydrocarbons. The presence ofcarbon dioxide in the products from the first pulse of ethylene, table 1, was directevidence for this. If we assume an oxygen monolayer on a surface containing1.1 x 1 0 l s atom/cm2, then the amount of carbon dioxide is within a factor of two ofthe calculated amount. No conclusions about the stoichiometry of the layer arepossible.Table 1 also shows, in agreement with previous work, that a considerable part ofthe ethylene which passed over the treated and untreated catalysts was retained.Thus rapid deactivation of the catalyst, table 1, can be accounted for in terms offormation of carbonaceous residues.These residues may well have covered theoriginal platinum grains and prevented their sintering. Fig. 8 can be interpreted interms of surface protection by deposits which inhibited sintering on subsequenttreatment by hydrogen.may be attributed to the slowbuild up of carbonaceous deposits rather than to sintering, for most of the latterprocess must already have taken place during pretreatment.In spite of the fact that the differences in particle size and surface area of thesecatalyst samples remained after they had been in contact with ethylene, the differencesin activity became negligible during subsequent ethylene pulses.This must havebeen caused by the rapid deactivation of the most active sites by hydrocarbon frag-Tompkins and AkhtarThe slow decrease in activity described by PaiT . BAIRD, z. PAAL AND s. J . THOMSON 55inents. This, in turn, suggests that only a fraction of the surface sites actually tookpart in the catalytic reactions.This sintering process under relatively very mild conditions emphasizes that thesystem Pt/H, may behave as a very mobile solid system and also the importance ofthe eventual support in keeping the individual catalyst granulates apart. On theother hand, the unusual activity of hydrogen in transforming the structure of theplatinum metal suggests once more that the hydrogen may play a more active rolethan had been suspected in creating active centres in platinum and other metals.The relationship between different pretreatments, excluding area effects, and catalyticactivity will form the subject of another study.The authors are very grateful to Dr.R. L. Moss of the Warren Spring Laboratory,Stevenage, for the X-ray investigations. One of us (Z. P.) thanks the I.A.E.A. for aFellowship.M. Boudart, A. W. Aldag, J. E. Benson, N. A. Dougharty and C. G. Harkins, J. Catalysis,1966, 6, 92.0. M. Poltorak and V. S. Boronin, Russ. J. Phys. Chem., 1966,40,1436.T. A. Dorling and R. L. Moss, J. Catalysis, 1966,5, 111. ' M. Boudart, A. W. Aldag, L. D. Ptak and J. E. Benson, J. Catalysis, 1968,11,35.Y . Barron, G. Maire, J. M. Muller and F. G. Gault, J. Catalysis, 1966, 5,428.G. Maire, C. Corolleur, C. Juttard and F. G. Gault, J. Catalysis, 1971,21,250. ' G. R. Wilson and W. K. Hall, J. Catalysis, 1970,17,190.G. R. Wilson and W. K. Hall, J. Catalysis, 1972, 24, 306.J. E. Benson and M. Boudart, J. Catalysis, 1965,4,704.l o D. E. Mears and R. C. Hansford, J. Catalysis, 1967,9, 125.l1 M. Akhtar and F. C. Tompkins, Tram. Farahy Soc., 1971,67,2454.l2 M. Akhtar and F. C . Tompkins, Trans. Fmday SOC., 1971,67,2461.l3 P. Steingaszner, T. Mhdy, Z. Schay and I. Kardos, Kiiolaj Fofdgcfz, 1968, 1, 151 ; Chern.l4 G. F. Taylor, S. J. Thomson and G. Webb, J. CataZysis, 1968,12, 191.l5 P. A. Sermon, J. Catalysis, 1972,24,460,467.l6 Z. Paill, Magyar Kbm. Folydirat, 1969,15,478.l7 P. TMnyi and L. Babernics, Acta Chim. Acad. Sci. Hungary, 1963,35,419.l8 G. F. Taylor, S. J. Thomson and G. Webb, J. Catalysis, 1968, 12, 150.l9 Z. Pa61 and S. J. Thomson, Radiochem. Radioanal. Letters, in press.Abs., 1969,71, 5098

ISSN:0300-9599    

DOI:10.1039/F19736900050

出版商:RSC    

年代:1973

数据来源: RSC

摘要:

Thermodynamic Behaviour of Acridine Orange in SolutionModel System for studying Stacking and Charge-effects on Self-aggregationBY B. H. ROBINSON,*? A. LOFFLER AND G. SCHWARZPhysical Chemistry Institute of the University of Basle, SwitzerlandReceived 3rd January, 1972The thermodynamics of association of the positively-charged dye Acridine Orange have beeninvestigated spectrophotometrically in water as a function of ionic strength and in the presence ofadded methanol, urea and dioxan. The extinction coefficients EM, ED, and Est, were measured whichare characteristic of the spectra of monomer, dimer and long aggregates of the dye molecules respec-tively. A theoretical model system, based on short-range stacking and long-range electrostaticinteractions, has been used to interpret the thermodynamic data.The results suggest that bothfactors are of importance in determining the tendency to aggregate. Furthermore, when the solventis varied they are related by a compensation effect ; decreases in stacking interactions are paralleledby decreases in repulsive electrostatic interactions.The results of this study on a model self-associating system are of relevance to the general problemof the importance of stacking and charge effects as factors determining ~ c e l l e formation, associationof nucleotides and nucleosides and the conformational stability of biological macromolecules insolution.Molecular aggregation in solution has been the subject of numerous investigationsin recent years.'. In biophysical chemistry, such phenomena are of particularinterest because of their significance with regard to the self-organisation of sub-unitsto large structures of biological activity.As model systems, the acridine series ofdyes may be employed. In this paper, we report a novel method of investigating thethermodynamics of association of acridine dyes in solution, which is based on spectralanalysis, and has wide general applicability. Additional information, e.g. the extentof interaction of the dye molecules in the associated complex, and the conformationof aggregates, can be derived from the shape and intensity of the spectra.3*Detailed results are reported for the association of Acridine Orange in solution,where it is shown that it is possible to evaluate, by empirical methods, the thermo-dynamics of the system over a wide concentration range. It is then possible toexamine in some detail the separate contribution of short range (stacking) and longrange (electrostatic) forces to the overall association.In a subsequent paper, we shall report results obtained by temperature-jumprelaxation spectroscopy on the kinetics and amplitudes of the relaxation processassociated with aggregati~n.~EXPERIMENTALThe Acridine Orange cation (fig.1) has been found to aggregate quite strongly in waterand so was suitable for detailed examination. The association was studied as a function ofionic strength and temperature in pure water, and in aqueous mixtures with methanol,dioxan, and urea. Solutions were buffered at pH = 6.9, using 1 mM Na2HP04+ KH,P04.Under these conditions, the dye is 100 % singly-protonated at the middle ring nitrogen(pK = 10.4 at 2OoQ6t present address : University Chemical Laboratory, Canterbury, Kent, England.5B.H. ROBINSON, A. L ~ F F L E R AND G. SCHWARZ 57AFIG. 1.-The Acridine Orange cation.The purchased sample of Acridine Orange hydrochloride (Fluka, Buchs, Switzerland)was grossly impure, and was initially purified by three recrystallisations from methanol.The crystals obtained were then redissolved in methanol+ water and 1 M sodium hydroxidewas added slowly with shaking. A yellow precipitate of the free dye formed, which wasfiltered, washed with water and vacuum dried. The melting-point was sharp at 182°C(lit. (Beilstein) 182"C), and the purity was confirmed by elemental analysis.Methanol, dioxan (Fluka puriss) and urea (Merck analysis) were used without furtherpurification.The ionic strength was generally controlled using sodium chloride (Flukapuriss). Sodium sulphate, calcium chloride and tetramethylammonium chloride were allMerck or Fluka puriss grade.Spectrophotometric measurements were made on a Cary 1605 spectrophotometer,equipped with a model 126 Recorder Interface. 1-cm, 0.1-cm and 0.01-cm cells were used,the temperature within the cells being controlled to +O.O2"C.CHARACTERISTICS OF THE ACRIDINE ORANGE (Go') SPECTRUMBeer's Law is not obeyed for the visible spectrum of the acridine series of dyes, andaggregation is accompanied by a progressive blue shift of the main absorption peak. Spec-tral analysis of overlapping absorption peaks in the presence of multiple equilibria is not easyand most analyses have been confined to consideration of only a monomer-dimer equili-brium '* * or of a particular monomer-n-mer eq~ilibrium.~For Acridine Orange aggregation, we find (fig.2) an approximate isosbestic point at lowconcentrations (1-5 x lob5 M) at 25°C in water, corresponding effectively to a monomer-dimerequilibrium over this concentration range. However, as the total concentration of AcridineOrange is further increased, it is clear that the shape of the spectrum is continually changing,and shows both a hypsochromic shift and hypochromism at the monomer peak wavelength.At any concentration, there will be up to three types of species present in solution, whichmay be defined as monomers, dimers, and oligomers, where oligomers are defined as aggre-gates of three or more Acridine Orange units.We may define the extinction coefficients ofthese species in terms of single Acridine Orange units, with reference only to the nearest-neighbour interaction of each unit. A monomer (unit) has solvent immediately on bothsides of the flat Acridine Orange molecule, and has an extinction coefficient defined by EM.If one side of the unit is occupied by an adjacent Acridine Orange unit, while the other side issolvent, we define the extinction coefficient as ED. If both immediately adjacent sites of theAcridine Orange unit are occupied by other Acridine Orange units, then we define theextinction coefficient of that unit as ESt.Furthermore, we will assume that cSt is independentof the length of the oligomer.The absorbance (E) of a solution of total concentration cx, in a cell of path-length I is thengiven by :E = E C ~ I = E~c~+~EDCC,+E~~C(~-~)C, ( I = 1 cm)2 2where E is the extinction coefficient of the solution and c, is the equilibrium concentration ofn-mer. Before the thermodynamics of association can be evaluated, it is clearly necessary toknow EM, ED, and ESt.DETERMINATION OF EM.-EM is given by the limiting extinction coefficient as cg is decreasedat 492 nm, and is in principle easily determined from the spectrum of a very dilute solutionof dye. However, in water as solvent at 25"C, Acridine Orange still contains a significan58amount of dimer even at lo-' M.Below this concentration, spectroscopic observationsbecome increasingly inaccurate because of adsorption of the charged dye on the walls of thecontaining flask and the optical cell, leading to low values of EM. By repeating experimentsin the same glassware, however, a reproducible and stable extinction coefficient can beobtained. The limiting extinction coefficient as cX-0 is 60 OOO& lo00 M-l cm-l at 492 nm,where 492 nm is the wavelength of maximum absorption of the monomer band. This valueof EM is found to be independent of ionic strength and temperature, and is in excellent agree-ment with other determinations in water.6* '* loTHERMODYNAMICS OF ACRIDINE ORANGE AGGREGATION!OFIG.2.-Concentration dependence of Acridine Orange spectrum : A, 5 x M ;C, 2.5 x M ; D, 5 x M ; E, 1 x M ; F, 5 x M. Temperature = 25T, I = 0.1 M(NaCl).M ; B, 1 xSince the equilibrium properties of the system were measured in aqueous media containingup to 10 % v/v methanol and dioxan, it is important to investigate the effect of changes insolvent composition on extinction coefficients. In 98 % v/v methanol, &hg2 = 72 0o0 M-lcm-l, but there is no significant shift in the absorption maximum. Similar differences arefound for transfer to other solvents. In view of these comparatively small overall effects,it can be assumed that E & ~ ~ is effectively unchanged by the addition of low concentrations(-10 %vlv) of additives to water, and this is confirmed experimentally.Slight batho-chromic shifts of the wavelength of maximum absorption of the monomer have been observedfor proflavin on transfer from water to ethanol-water mixtures,l while small hypsochromicshifts are observed for Methylene Blue in the same solvent media.12 These differences inbehaviour of EM for the different dyes presumably reflect the smdl changes in dye structure.DETERMINATION OF ED.-AS the concentration of Acridine Orange in water is increased,it is apparent that there is no pronounced isosbestic point over any wide concentration range.Therefore, to determine the spectrum of the dimer, it is necessary to work at low totalconcentrations (less than M) where onIy monomer and dimer are present at significantconcentrations, and assume only a monomer-dimer equilibrium over this range.Adsorptionof dye on glassware and optical cells must be prevented to obtain accurate results. Fromwork done under these conditions, values of 15 500 M-l ~ m - l , ~ 16 OOO M-l cm-l and13 O00 M-' cm-l l4 have been obtained for ED at 492 nm. (Considerably lower values of EB . H. ROBINSON, A . LOFFLER AND G . SCHWARZ 59are obtained when higher concentrations (> M) of Acridine Orange are used to evaluate~$92 on the basis of a monomer-dimer equilibrium.) From the figures quoted above, a valueof E $ : ~ of 15 OOO M-' cm-l may be arbitrarily chosen for use in the calculations to bedescribed later, since the thermodynamic results are found not to be critically dependent onthe values chosen for the extinction coefficients.M), the solutiontends increasingly to exhibit a spectrum characteristic of stacked molecules, i.e.longeraggregates are formed at high concentrations. This tendency to aggregation is furtherpromoted by increasing the ionic strength or decreasing the temperature.In principle, it is possible to measure a limiting spectrum at high concentrations which isessentially the spectrum of long aggregates, with extinction coefficient ESt. However, evenworking with 0.01 cm cells and high values of optical density, the spectrum is still exhibitingsmall changes with increasing cx. To obtain a limiting value of &St, an extrapolation pro-cedure must be applied.An exact extrapolation procedure depends on a detailed prior knowledge of the thermo-dynamics of the aggregation process, and so we are forced to use empirical methods ofevaluation.A rigorous approach is not required as we are concerned with only a smallextrapolation. A suitable procedure was found to be the zero extrapolation of a plot of Eagainst Ki$ (at a fixed high value of c i , and variable temperature), where Klz is themonomer-dimer equilibrium constant measured at low concentrations. Values of Klat other temperatures were obtained using the experimentally determined value of AH,", =- 38 kJ rn~l-'.~ These plots were found to be linear, and only a small extrapolation wasrequired to K;$ = 0 (which implies maximum possible stacking for a reasonable tendency toaggregate), when E = &St (fig.3). If strong aggregation is assumed as in Model 2 (to bedescribed later), then, for (Kl2ci)% 1, we find :-DETERMINATION OF &st.-As more Acridine Orange is added/ o-o-o0-0 LdOU20 ;c0 1 2 3 4 5 6K;:/moll.-'FIG. 3.-Plot for evaluation of ESt at fixed c;, and variable temperature.Since, in addition, &~/(Kl2c~)<2E~/(Klzci)3, a plot of &a against K T t should be linear withintercept &St. Linear plots were indeed obtained, with ESt indistinguishable from that derivedfrom the plot of E against Ki; described above.A value of ~ & 9 ~ of 6500 M-l cm-' was derived from fig. 3 and is used in all subsequentcalculations. Supporting evidence for this derived value is that it is identical with the valueof ~ t ? ~ determined for Acridine Orange stacked adjacent to the glutamate residues of the coilform of poly(g1utamic acid) at pH = 6.9.14 In addition ESt is independent of ionic strengthin this case.The agreement between these two separate determinations suggests that EStis not significantly dependent on environment. We will therefore assume that ESt is inde-pendent of ionic strength, small variations in solvent composition and aggregate length.I5From the plots in fig. 3, measured at different wavelengths, the spectrum of a stacked dyeunit in solution is readily obtained60 THERMODYNAMICS OF ACRIOINE ORANGE AGGREGATIONThe visible spectra of monomer, dimer (taken from ref. (7)) and long aggregates are shownin fig. 4.I Ah/nmFIG. 4.-Spectra of monomer, dimer and stack of Acridine Orange in solution : A, monomer spec-trum ; B, dimer spectrum ; C, stacked spectrum.THEORYTHEORETICAL BASIS FOR THE AGGREGATION PROCESS AND METHOD OFCALCULATION OF THE CORRESPONDING EQUILIBRIUM CONSTANTS FORASSOCIATIONAfter evaluation of the necessary spectroscopic quantities, we can now considerquantitatively the thermodynamics of the association process.The equilibriainvolved may be expressed as :Kn,n + 1Cn+C1 + Cn+1 (n = 1+c0).A special case of this equilibrium system is one in which the equilibria display " all ornone " characteristics, in which case the process may be simply considered as :K1 rn nc, + cn,where the particular aggregate, c,, is especially favoured, as in the case of micelleformation. Such aggregation processes are characterised by abrupt changes inobservable parameters at the critical micelle concentration. However, no such sharptransitions are observed in the spectrum of Acridine Orange as a function of concentra-tion.In this aggregating system where the Acridine Orange units are thought to bestacked in a sandwich arrangement at a small angle to each other,' it would seemdifficult to justify energetically the preferred existence of one particular aggregate,and so the micelle model will not be considered further. The micelle case is, however,readily analysed, since one equilibrium constant Kl ,n predominates. The generalcase is much more complex since (n - 1) equilibrium constants must be evaluated for asystem with (n) significant species present.Clearly, it is impossible to evaluate allequilibrium constants Kn,n+l (n = l - . ~ ) , unless they are assumed to be related insome way. As a theoretical basis for discussing the aggregation process, we havB . H. ROBINSON, A . LGFFLER AND G. SCHWARZ 61developed an approach based on the existence of short-range stacking interactions(including hydrophobic and dispersive forces) and long-range electrostatic interactions(but neglecting specific solvent effects favouring a particular n-mer) which is consistentwith the assumption of an inter-relationship between the Kn,n+l values. The expectedbehaviour of the system is then evaluated for three cases of particular interest.Fig. 5 shows the fundamental process of aggregation of molecular units, from astack of length (n) units to one of (n+ 1) units.The equilibrium constant is givenby :Kn,n+ = c,+ /cnc, (neglecting activity coefficients): a :FIG. 5.-Schematic model for Acridine Orange aggregation.and the corresponding standard free energy change is :where AGE, represents the contribution to the overall standard free energy changefrom short range stacking interactions (independent of ‘‘ n ”) and AG$,, is a coulombicterm (depending on n) representing the summation of all the electrostatic interactionterms (based on the assumption of localised charge on each acridine orange unit),given by :In eqn (2), NA is Avogadro’s constant, eo the electronic charge, a is the averageseparation between Acridine Orange units in the stack, and Q is a screening factor,which is equal to the effective dielectric constant experienced by the stack.As Qis increased, the repulsive effect of charges in the stack is correspondingly diminished,i.e. the charges are more effectively screened from one another. Q can vary between1 and co.It should be emphasised that changes in Q are related to the electrostatic interactionwhich essentially controls the activity coefficients fn of the aggregates.The summation term represents the electrostatic repulsion potential of the addedacridine orange unit with the first, second, . . . nth, etc. nearest neighbour unit in thestack.and from eqn (1) and (2) := exp( - AG&/RT)exp - q z ( l / n ) { r }whereq = NAei/aaRT.Letexp { - AG,“,/RT) = K62ThenForwhereTHERMODYNAMICS OF ACRIDINE ORANGE AGGREGATIONn)} = c = 0.57721 (Euler’s constant).n2n + a , expC(l/n) = n/b(3)(4)b = exp (-(l-c)> = 1.53.For smaller values of n, b is no longer constant, and varies from 1.33 at n = 2 to 1.53at n = co.However, a satisfactory approximation is made by assuming b to be aconstant equal to 1.4. Then from (2) and (4), the theory gives :Kn,n+l 2: Kl,(bln)4 < K12 (rt = 2, 3, . . . a). ( 5 )Let us now consider three special cases of eqn (5).CASE l.-strong electrostatic repulsion ; i.e. 0-+1 or q91. Then :K12%Kn,n+l (n = 2-+a).This is the usual assumption made when studying dye association, when only K12 isconsidered significant.CASE 2.-Weak electrostatic repulsion ; i.e.04a or 4.4 1. In the limit, o = 00and there is no electrostatic repulsion or alternatively complete shielding. Then :K12 = Kn,n+1 (n = 2 + ~ ) ,i.e. all the equilibrium constants are equal.CASE 3.-Medium electrostatic repulsion ; i.e.o = (N,eZ/aRT) or q = 1.Then :Kn,n4 = Kl,(b/n) (ti = 2 3 ~ 0 ) .The equilibrium constants Kn,n ,- slowly decrease as n is increased.For the detailed analysis of these three cases, a dimensionless term s is introduced,which is always equal to K12c1, where c1 is the equilibrium concentration of monomer.This device greatly simplifies the calculations, and enables extinction coefficients andconcentrations of individual species (c2, c3, . . . c,) to be expressed as a function of s ina general way.DETAILS OF CASE l.-Kn,n+l = 0 (n = 2- 00).Total optical density of the system(I = 1 cm) is:E = E C ~ = E M C ~ + ~ E D c ~ ( 6 B. 14. ROBINSON, A . LOFFLER A N D G. SCHWARZThe equation for mass conservation is :c; = c1+2c2.The mass action condition is :cz = CIS.Eqii (6) and (7) give :E = EM( 1 + 2 J . y + r,2s( 1 + 2 s y .Froin (7) and (8), we obtain :K12 = ~ ( l +~S)/C;.General points for cases 2 and 3 : total optical density ( I = 1 cm) is :E = EC; = EMC1+ 2ED& + E&(n - 2)c,.2 2Mass conservation gives :C; = c,+Cnc, = ~,+2'&~+C(n--2)c,.2 2 2DETAILS OF CASE 2.-K1, = Kn,n,-l. The mass action condition is :c,+1 = CIS"From (13) :andFrom (111, (12), (141, (15):E = EM( 1 - s)2 + ED24 1 - s) + &&andK 1 2 = s/(l - s ) ~ c ~ .DETAILS OF CASE 3.-Kn,,,k1 = K,,(b/n) (n = 2-m).The mass action conditioiiis :c , , + ~ = (c,/b)(bs)"[n! n = 1-m= (c1/b2)~/~s(6s(exp (6s) - 1)). (1964 THERMODYNAMICS OF ACRIDINE ORANGE AGGREGATIONUsing (ll), (12), (18)+(19):andKI2 = s((l+bs) exp (bs)+(b-l))/bcl. (21)It should be noted that, for all three cases, E is dependent only on s, b, and the valueschosen for E ~ , cD and &st. The first procedure, therefore, is to calculate theoreticalcurves for E as a function of s for the various models.The profile of the resulting curve is fixed for a given system, and is only dependenton the choice of wavelength to be used for evaluation of the data. From the measuredextinction coefficient of a solution of given ci, we may then immediately determine sgraphically, according to which model is used.Since K12 is dependent only on s, c iand b (for case 3), it is calculated directly from eqn (lo), (17) and (21). Once K , , isestablished, all the Kn,,+l values are immediately known for a given model. The casewhich best describes the association process will be that which gives the most constantK12 over the whole concentration range.In fig. 6, calculated plots of E against s for cases 1, 2 and 3 respectively are shownat 1 = 492 nm (the monomer peak). For case 1, s can take any value between 0 andGO ; for case 2, the limiting value of " s ", as ci-+ co, is unity. For case 3, s canexhibit values from 0 to a practical limit of s 2: 10.I0 0.2 O f , 06 0 8 10 12 1 4 16 I F 2 0SFIG.6.-Dependence of E on s : 0, case 1 ; 0, case 3 ; A, case 2.RESULTSAcridine Orange aggregation has been investigated in a variety of mixed aqueoussolvent media and at several ionic strengths using different salts.METHOD OF ANALYSISTo demonstrate the method of analysis, we may use results obtained in water atI = 0.1 and 0 (table 1). At I = 0.1 and for case 3 (q = l), the K , , values show noobservable trend with increasing ci, but just random scatter, reflecting experimentalerrors in the determination of 8. There are marked trends if results are evaluated onthe basis of the other two models (cases 1 and 2) with deviations in opposite directionsB . H. ROBINSON, A . LOFFLER AND G. SCHWARZ 65TABLE AP APPARENT K12 VALUES AS A FUNCTION OF CONCENTRATION FOR ACRIDINE ORANGEAGGREGATION IN WATER AT 25'CI = 0.1 M (NaCl)10-4 K121M-1case 1 case 2 case 3I = O10-4 ElpIM-1case 1 case 2 case 31.66 2.02 1.54 1.69 1.08 0.88 0.974.16 2.62 1.38 1.69 1.05 0.69 0.798.33 2.92 1.18 1.53 1.06 0.60 0.7116.6 4.41 1.06 1.55 1.20 0.51 0.6641.6 - 0.93 1.72 1.24 0.39 0.57limiting value, &+O; K12 = 1 .6 ~ lo4 M-IFor case 1, at I = 0.1, the analysis reveals an apparent increase in K12 with increasingc i , and hence the tendency to association is underestimated. When the modeloverestimates the aggregation tendency (case 2), this is revealed by apparently de-creasing values of K12 as c i is increased. It is clear that case 3 gives a best-fit to thedata at I = 0.1, and this therefore implies that the Kn,n+l values slowly decrease as nis increased.At I = 0, case 1 gives a better fit than case 3, and so the aggregationbehaviour can best be explained with a value of q between 1 and a. This impliesthat the tendency to aggregate beyond the dimer stage is not very strong at I = 0,and that the screening factor 0 increases as the ionic strength is increased. The exactvalue of K,, for any model is measured as cx-0, and may be evaluated by any of threepossible extrapolation procedures, the one employed in a given case depending onwhich model best approximates the aggregation process. These plots are shown intable 2. However, the analysis we have used does not justify a detailed analysis of r~(or q) values, so we have assigned CT semi-quantitatively to one of five possible ranges.case 1 qa 1 0-blintermediate between case 1 and case 3 q > l l<o<NAe;/aRTcase 3 q = 1 0 = NAe;/aRTintermediate between case 3 and case 2 NAe3 /aRT< Q < cocase 2 4 - a 6+00K12 = 1 .0 ~ lo4 M-'q < lTABLE 2.-PLOTS FOR THE DETERMINATION OF Ki2bat-fitmodel Y x grad. int.1 cxls 2s KiJ k122 CXlS s(s- 2) K J Kii?3 ci/s (1 + bs)ebs/b G - 2 l (b- 1)/bKi$EFFECT OF ADDITION OF SODIUM CHLORIDE ON K12 AND QThe ionic strength dependence (NaCl) of K,, and q (in water at 25°C) is shown intable 3. It is found that K12 steadily increases as the ionic strength is raised, withlog K12 proportional to ,/I, as predicted by Debye-Hiickel theory. The q valuesdecrease (0 values increase) quite sharply up to I = 0.1, when q 21 1 and then levelOff.TABLE IO IONIC STRENGTH DEPENDENCE (NaCl) OF K12 AND q IN WATER AT 25OC0 1 .o $ 10.01 1.2 >10.05 1.4 >10.1 1.6 10.25 2.2 10.50 3.2 1I t M 10-4 K1pIM-1 41-66 THERMODYNAMICS OF ACRIDINE ORANGE AGGREGATIONEFFECT OF SOLVENT ADDITIVES ON K12 AND 4The effect of addition of methanol, urea and dioxan at I = 0 and 0.1 is shown intable 4.The effective dielectric constant 0 is slightly higher in 10 % methanolcompared with pure water, and there is a progressive decrease of K,, in 5 % and 10 %methanol. The data at I = 0. I , and 10 % dioxan suggest that the effect of the chargesTABLE 4.-EFFECT OF ADDITION OF METHANOL, DIOXAN AND UREA AT I = 0 AND 0.1 M (NaCI)molefraction I = O I = 0.1 Madditive (Nsdd) vol.% 10-4 K ~ ~ / M - ~ 4 10-4 K ~ ~ I M - I 4methanol 0 0 1 .o $ 1 1.6 10.023 5 0.55 >1 1.1 10.047 1 0 0.33 1 0.73 1dioxan 0 0 1 .o 1.6 10.01 1 5 0.2 1 0.5 < 10.023 1 0 0.08 1 0.14 0urea 0.035 molarity 0.25 > I 0.55 12Mon the dye units is almost totally neutralized (case 2 obeyed). Dioxan also producesthe largest proportional decrease in K12 of the three additives. Urea produces amarked decrease of K12, and a moderate increase in the aggregation tendency, com-pared to the pure water values. We may conclude from the results that for any ionicstrength I, the order of CT values is :and this is paralleled by the order of Kl, values :water < methanol c urea c dioxanwater > methanol > urea > dioxan.EFFECT OF OTHER ADDED SALTSThe effect of sodium sulphate, calcium chloride and tetramethylammoniumchloride on the K12 values is shown in fig.7. On an ionic strength basis, the order ofincreasing K12 values isNaCl> CaCl, > Na,SO,.Addition of tetramethylammonium chloride shows a maximum value of KI2 atI = 0.5. The q values suggest that case 3 (q = 1) gives a good fit to the data over awide range of ionic strengths, although at high values of ionic strength for calciumchloride (I = 0.75 and 1.5) and tetramethylammonium chloride (I = 1 .O), q becomesless than unity, and stacking is further promoted.DISCUSSIONThe visible spectrum of Acridine Orange is interpreted as showing the formationof dimers and oligomers as the total concentration of dye is increased.From thespectra, measured as a function of concentration, EM, E~ and ESt can be evaluated.Using these experimentally derived values, the equilibrium constants characterisingthe aggregation process can be calculated on the basis of several possible models.We have assumed the two principal factors determining aggregation to be short-rangestacking interactions, AG&, and long range electrostatic interactions, AGgl. Further-more, we have assumed that the geometry of the associated species, and hence cM, cDand Est are unchanged with ionic strength or small additions of co-solvent, and that &SB. H. ROBINSON, A . LOFFLER AND G . SCHWARZ 61is independent of stack length. It was found possible to evaluate the detailed equi-libria on the basis of the above theoretical model without recourse to more complexequilibrium behaviour (involving second nearest neighbour stacking interactions),which was found to be necessary in the case of Methylene Blue aggregation in water.'4.0// ii /i /// Lot , , , *0 0.4 0.8 1.2 1.6IIMFIG.7.-Effect of ionic strength on K12 using various salts at 250°C : (- . - . - 0), NaCl ;(---- +), CaC12 ; (- *), Na2S04 ; (- . . - . . - x ), (CH&NCl.ANALYSIS OF THE SCREENING FACTOREFFECT OF IONIC STRENGTH ON QIncrease of ionic strength for the solvent systems we have studied always gives anincrease in Q, or an increase in the tendency to aggregation. This behaviour can beunderstood in terms of the screening of positive charges in the aggregate through theattraction of negative charges from the surrounding medium, resulting in a negative" charge cloud " around the Acridine Orange aggregate.At Z = 0, the tendency toaggregate is not pronounced since the measured screening factor is small. For manyvalues of the ionic strength, q N 1, and a direct determination of the absolute value ofQ is then possible, if the separation of dye units in the stack is known. Since q = 1,Q = iVei/aRT. If we take a = 5.5A (which is an average value between thatobtained in solution from spectral considerations 4* and by X-ray crystallography16)and make the assumption of localised charge on the Acridine Orange units, thenQ ~li 100. This value may be thought to be equivalent to the effective dielectricconstant experienced by the microscopic system.EFFECT OF CO-SOLVENTS ON 0Addition of methanol, dioxan and urea also increase 6, which is perhaps a sur-prising result, since the bulk dielectric constants of the additives are less than for water.Dioxan gives especially pronounced effects.An explanation of this effect requires adetailed knowledge of solvation around ions, but a tentative explanation may beassociated with an increase in the overall local dipole moment around the stack onaddition of co-solvent. The results demonstrate clearly that it is invalid to use bulkdielectric constant values and properties to interactions in microscopic environmentsof the solvent system68 THERMODYNAMlCS OF ACRIDINE ORANGE AGGREGATIONANALYSIS OF K12 VALUESIt has already been shown (eqn (1)) that K12 is a quantity comprising both stackingandelectrostaticcontributions.The stacking contribution itself consists of three factors,These are : mixing effect (mix) ; dispersive interactions (disp) ; hydrophobicforces (hyd). ThereforeThese contributions to AG;12 may to some extent be treated separately, following theprocedure of Mukerjee.2AGhi,.-Measured equilibrium constants have been expressed in niolarity units.To correct for the change in the number of solute particles through the reaction, it isnecessary to express Kin mole fraction units. This contribution, for converting twomonomer species into a dimer species, gives AG& = +9.9 kJ mol-l,17 independentof solvent, temperature, and ionic strength.AGgl.-Since AG,"I = iVei/ac for dimer formation, and Q is always positive, thisterm will lead to an increase in K12 as CT is increased.When q = 1, as it is under manyof the experimental conditions we have employed, then AGZ1 is +2.5 kJ mol-l. Inpure water, at zero ionic strength, the electrostatic contribution is considerablygreater, but when case 2 is obeyed (e.g. 10 % dioxan I = 0.1) then AG,"' = 0.AG&,,.--Direct calculation of the dispersive interactions involved in the acridineorange system is not possible because of the complex structures involved but attractiveforces through n-n interactions in the dimer will result in an overall negative AG&.Experimental approaches to AG& based on interfacial free energy considerations donot clearly separate the contributions made by AGiisp and diGg',d.AGiyd.-AGgyd is composed of interactions resulting from changes in solvation.Around large molecules and ions in water, the region of solvent immediately adjacentto the solute is more structured 1 8 p 2 * than the bulk solvent.Consequently the entropyof a water molecule in the solvation shell of the Acridine Orange monomer is lower thanin the bulk solvent. It the temperature is increased, or a structure-breaker is added,the solvation shell is progressively broken down. On formation of a dimer from twomonomer molecules, two faces of Acridine Orange molecules in contact with water inthe solvation shell are lost, and two regions of structured water are changed intobulk solvent molecules with less structure and higher entropy (i.e.AS&,> 0). It isto be expected that increasing the librational motions of water molecules will alsorequire compensatory increases in the enthalpy of the system (i.e. AHCyyd>O), but theoverall free energy change AGiyd will be negative. As for AGiisp, it is not easilypossible to evaluate the separate contribution of AG&d to AG;I2.However, the combined effect of AG& and AG& lo the overall stability of thedimer may be estimated. To take a general example, that of K12 in 0.1 M NaCl, wemeasure AGY2 N -24 kJ mol-l. Substituting AG&, = 10 kJ mol-1 and AG,", =2.5 kJ mol-l in eqn (22), gives a value of AG;: (where AG&' is AG& corrected formixing) equal to -36.5 kJ mol-'.Since AHf2 is -38 kJ r n ~ l - l , ~ the entropy change on dimerisation (corrected formixing) is slightly negative, and this would appear to suggest that hydrophobicinteractions do not make the major contribution to AGE,, whereas it is implied thatAH&, is highly negative and probably greater than -40 kJ mol-1 because all otherenthalpy contributions are likely to be positive.Therefore, if the above thermodyna-mic analysis of hydrophobic forces is correct,21 dispersive interactions are the majorcontributor to the stability of the Acridine Orange dimer. However, for micelle for-mation by surface-active agents, TAS"> with AHo approximately zero, and thissuggests that hydrophobic interactions provide the main driving-force for micellization.AGY2 = AG~i,+AG,"+AG~;,,+AG~,.d. (22B .H . ROBINSON, A. LOFFLER AND G. SCHWARZ 69EFFECT OF IONIC STRENGTH ON K12A major effect of the added salt will be caused by the electrostatic interactionmentioned previously in connection with the stabilisation of oligomers, where adecrease in repulsion between the charges in the stack is expected due to the preferredattraction of anion from the solution. This hypothesis is supported by the muchgreater increase in K1 obtained using calcium chloride rather than sodium sulphateat equivalent ionic strengths. The results therefore suggest that an effect on AG:,through changing c is the major cause of the increase in K , , with ionic strength.However, addition of salts will also cause some modifications in solvent structurewhich may contribute to a smaller extent to the overall change in K I 2 .The resultsobtained with tetramethylammonium chloride show this effect more clearly. At lowionic strengths Z<0.2, there is a pronounced increase in KI2 due to the electrostaticshielding effect, but K I 2 levels off around I = 0.5. On further addition of tetra-methylammonium chloride above Z = 0.5, K1 decreases again, presumably becauseAC& becomes less negative, due to a solvent structure-breaking effect of the tetra-met h ylammonium cat ion.EFFECT OF CO-SOLVENTS ON K12Addition of methanol, urea, and dioxan shows the operation of hydrophobicforces in the system. For a given ionic strength the order of K l z values (cal-culated on a mole fraction basis) is water > methanol> urea > dioxan. The re-verse order represents the relative structure breaking efficiency of the additives.We find that dioxan, in addition to being the best structure breaker, is also the mostefficient additive for screening the charges in the aggregate, and therefore, thestructure-breaking and charge screening roles of the additives compensate each otherto some extent in the process of self-aggregation of charged species.The authors thank Dr. V. Vitagliano, Dr. W. Balthasar and Dr. J. Seelig forhelpful discussions. B. H. R. thanks the Royal Society of Great Britain for aEuropean Fellowship.Molecular Association in Biology, ed. B. Pullman (Academic, New York, 1968).P. Mukerjee and A. K. Ghosh, J. Amer. Chem. Soc., 1970, 92, 6403.H. DeVoe, J. Chem. Phys., 1964,41, 393.R. E, Ballard and C. H. Park, J. Chem. SOC. A, 1970, 1340.B. H. Robinson, A. Loffler and G. Schwarz, to be published. ' V. Zanker, Z. phys. Chem., 1952,199,225. ' M. E. Lamm and D. M. Neville Jr., J. Phys. Chem., 1965, 69, 3872.G. Schwarz and W. Balthasar, European J . Biochem., 1970,12,454.M. Hida and T. Sanuki, Bull. Chem. SOC. Japan, 1970,43,2291.l o R. W. Armstrong, T. Kurucsev and U. P. Strauss, J. Amer. Chem. SOC., 1970,92,3174. '' G. Lober, in Interactionen bei Biopolymeren (Studia Biophysica, Berlin, Vol. 24/25, 1970), p.l 2 P. Mukerjee and A. K. Ghosh, J. Phys. Chem., 1963,67,193.l 3 F . W. Schneider, S. K. Podder and M. Eigen, unpublished results.l 4 B. H. Robinson, A. Loffler and G. Schwarz, to be published.l 5 D. J. Blears and S. S. Danylik, J. Amer. Chem. SOC., 1966, 88, 1084.l 6 P. J. Wheatley, J. Chem. SOC., 1959, 3245.l 7 R. W. Gurney, Ionic Processes in Solution (McGraw-Hill, New York, 1953).238.For a detailed discussion of the meaning of the term " water-structure " see Symposium on H-bonded Solvent Systems, ed. A. K. Covington and P. Jones (Taylor and Francis, London, 1968),p. 224-6.l9 H. S. Frank and M. W. Evans, J. Chem. Phys., 1945,13,50.2o G . NCmethy and H. Scheraga, J. Chem. Phys., 1962,36,3401.21 A. Holtzer and M. F. Emerson, J. Phys. Chem., 1969,73,26.22 M . F. Emerson and A. Holtzer, J. Phys. Chem., 1967, 71, 3320

ISSN:0300-9599    

DOI:10.1039/F19736900056

出版商:RSC    

年代:1973

数据来源: RSC

摘要:

Prolonged Irradiation of Dilute Aqueous Solutions ofRedox CouplesBY P. L. AIREYIsotope Division, Australian Atomic Energy Commission,Lucas Heights, N.S.W., AustraliaReceived, 24th January, 1972On prolonged irradiation, the potential of an inert electrode immersed in aqueous solutions of awide range of redox couples assumes a steady value (equivalent redox potential, ERP) which isdetermined in some cases by hydrogen peroxide and in others by the redox couple. Irradiation ofarbitrary concentrations of these redox couples leads to net oxidation or reduction according towhether the ERP is greater or less than the equilibrium potential of the couple. A theory is proposedwhich defines the conditions under which the ERP is established and predicts that when it is redoxcon trolled,ERP = E(HzOz)-Awhere A is 24 mV for T o y-radiation and - 1.8 mV for ‘OPo =-particles (0 < pH < 2).The theory predicts that the ERP is independent of the nature of the redox couple and of theintensity of the radiation ; it also explains why the ERP is not established on prolonged irradiation ofredox dyes.1.INTRODUCTIONProlonged irradiation of dilute aqueous solutions of a wide range of redox couplesleads to a stationary state in which the concentration ratio of the componentsdepends on the relative velocities of reactions with the following species formed in theprimary radiation step :(1 * 1)Dainton and Collinson pointed out that the nature of the stationary state could bepredicted from the standard electrode potential of the couple.From a survey of theliterature (1950) they concluded that X-irradiation would lead to complete oxidationif Eo were less than 0.52V (I.U.P.A.C. Stockholm Convention) and to completereduction if Eo were greater than 1.1 V. If Eo lay within this range, the stationarystate would comprise appreciable concentrations of both components. They there-fore proposed that irradiated aqueous solutions were characterized by a potentialwhich they called the equivalent redox potential (ERP). Irradiation of diluteaqueous solutions of redox couples would induce net oxidation or reduction accordingto whether the ERP were greater or less than the equilibrium potential of the couple.From the data available at the time, a value of f0.95 V was estimated for the ERP.Dainton and Collinson commented that this value could not be explained simply interms of the reactions of H- atoms and OH= radicals with the components of the redoxcouples.They also pointed out that the observed reduction of aqueous solutions ofthe redox dyes Methylene Blue, indophenol, phenosafranin, resazurin and triphenyl-tetrazolium chloride (- 0.35 < E,/V< 0.45) is not compatible with a value for theERP of 0.95 V. proposed that the ERP was a kinetic potential depend-ing on the rates of the reactions of the radiolytic products at the electrode surface.H20-H2, H202, He, .OH, e& H+, OH-.Veselovskii7P. L . AIREY 71The first accurate measurements were performed by Henderson et aL4 whoconcluded from studies of the prolonged 6oCo-y irradiation of a wide range of redoxcouples (see table 1) thatCartledge suggested that the principal radiation electrochemical effects were due tothe stable radiolytic products and that the effect of the redox couple was to stabilisethe potential.The measurements of Henderson et al. were confirmed over the wholepH range by Feates and Knight who pointed out that, within experimental error, theERP was numerically equal to the hydrogen peroxide potential E(H202), which wasdetermined by Bockris and Oldfield ' to beERP = 0.85+0.02-0.059 pH V (O<pH< 14). (1 -2)E(H202) = 0.835 & 0.005 - 0.059 pH V (0 < pH < 14). (1 -3)Feates and Knight assumed that H202 was the potential determining species in theirexperiments and suggested that the results of Henderson et al.might be interpretedsimilarly.In this paper, explanations are offered for the following properties of the ERP onthe basis of electrochemical and radiation chemical kinetic considerations. (i) Evenwhen the redox couple is potential determining, the ERP is numerically very close toE(H202). (ii) It is independent of the nature of the couple and the dose rate, andalmost independent of the radiation quality. (iii) The ERP is observed even ifstandard potentials of the redox couples are far removed from the ERP. (iv) Thebehaviour of irradiated aqueous solutions of redox dyestuffs cannot be predicted fromthe ERP and the equilibrium potentials.The argument is developed along the following lines.Methods for establishingthe nature of the potential controlling species are outlined in section 2. The value ofthe ERP is calculated in two stages. In section 3(i), it is shown that under specifiedconditions, the potential reached on continuous slow addition of hydrogen peroxide todilute aqueous solutions of a redox couple is equal to E(H202). Kinetic implicationsare discussed in section 3(ii). As a result of radical reactions the ERP differs slightlyfrom E(H202). The magnitudes of the differences, which are equal to the postirradiation effects, are calculated in section 3(iii). Finally applications and generalis-ations of the theory are considered in sections 4 and 5.2. THE NATURE OF THE POTENTIAL DETERMINING SPECIESOn prolonged irradiation of solutions of redox couples the potential of a platinumelectrode could be controlled by electrochemical reactions involving H2, H202(eqn (1.1)) or the redox couple.Since the ERP is close to E(H202), electrochemicaloxidation of the molecular hydrogen cannot be kinetically significant.Whether H202 or the redox couple is potential determining depends on the ratioof the parameters io (redox)/co(H202). The parameter io (redox) is the exchangecurrent density associated withThe H202 potential is a mixed potential invoIving the following reactions+ e,+M"+. (2.1)2H+ + 0, +2e&H,02, Eo = 0.682 V (2.2)2H++H,02+2e,f2H20, E,, = 1.771 V. (2.3)M(n+l)fAt the hydrogen peroxide potential there is no net current flow. Thus the partialoxidation and reduction current densities are equal and are denoted by c0(H20,)72 EQUIVALENT REDOX POTENTIALThe following criterion is used to decide which species controls the potential in theirradiated solution. Ifthe redox couple will be potential determining, whereas ifH202 will control the potential.activities of the reactants.andwhere a is the transfer coefficient, and at120z, a.and aR are the activities of H202,M("+l)+ and M"+ respectively. The parameters cOO(H202) and ioo (redox) aredehed by eqn (2.5) and (2.6).The steady state activities of the redox components at the ERP are determinedfrom the total analytic concentration, the activity coefficients and the activity ratiowhich is calculated from the Nernst equation :Of the systems listed in table 1, only in the ferrous-ferric and iodide-iodine cases areappreciable concentrations of both components present on prolonged irradiation.In these cases the hydrogen peroxide activity must be known before the nature of thepotential determining species can be decided.io(redox) 2 10 cO(H202) (2.44c0(H202) b 10 io(redox) (2.4b)Both co(H,02) * and i,(redox) are functions of theco(H202) = C O O ( H 2 O 2 ) ~ H 2 O r (2.5)io(redox) = ioo(redox)aia&' - a ) (2.6)E = ERP = E,(redox)+(RT/F) In (ao/aR). (2.7)totalconc.1moI 1.-110-410-310-32 .5 ~ 10-45 . o ~ 10-45 . o ~ 10-45 . o ~ 10-48 . o ~ 10-410-410-35.0 x 10-410-310-310-410-310-310-46 . o ~ 10-410-3TABLE 1.ERP/V(w. H2elect rodePH at same pH)0.50 0.83-0.13 0.840 0.840.50 0.870.90 0.852.09 0.869.47 0.8411.35 0.851.4 M NaOH 0.880.50 0.870.0 0.850.37 0.8113.9 0.894.0 0.840.37 0.860.90 0.850.50 0.836.4 0.8766.4 0.87513.8 0.835EOP0.6*0.60.60.60.60.6g60.691.711.711.711.711.250.290.23potentialdeterminingspecies ref.H202 4redox 4redox 6redox 4redox 4redox 4?p 4? 4? 4? 4? 6redox 4redox 6H202 4H202 4H202 6H202 6H202 4H201 4H202 4* Because of the equilibrium concentration of I; which is allowed for in deciding the nature of thepotential determining species, the standard electrode potential Eo lies in the rangeEo(Ii-I-)(= 0.54 V)<Eo<E(Iz,-I-)(= 0.63 V).7 Insufficient kinetic data is available to permit an assessmentP.L. AIREY 73The steady state hydrogen peroxide concentration can be calculated by assumingthat it is generated radiolytically and decomposed by thermal reactions with the redoxcomponents. The concentrations of the added solutes are such that reactions of theradiolytic intermediates (He and OH. at pH<3) with the hydrogen peroxide can beneglected.2Mn+ + H202 + 2H++2M("+ l)+ + 2H202M'"+ l)+ + HzOz +2Mn+ + 2Hf + O2daHzO,/dt = GH20zD'- v2.8- v2.9(2.8)(2.9)(2.10)Thuswhere G H ~ ~ ~ is the radiation chemical yield of HaOz, D is the dose rate (eV 1.-' s-l),D' = 0/1000N and N is Avogadro's constant. v 2 . 8 and V2.9 are the velocities(moll.-l s-l) of reactions (2.8) and (2.9) and are, in general, given byand(2.1 1)(2.12)where k2.8 and k2.9 are rate constants andf and g are functions determined by thereaction kinetics.The steady state hydrogen peroxide activity is found from eqn(2.10), (2.11) and (2.12).Having calculated ao, aR and aHzo2 under conditions prevailing at the ERP, theparameters co(H202) and &,(redox) can be deduced from the literature values usingeqn (2.5) and (2.6). The nature of the potential determining species is decided fromthe relative magnitudes of the parameters (eqn (2.4a) and (2.4b)). Typical calculationsare outlined in Appendix I. The nature of the potential determining species so foundhave been listed in table 1 for those systems in which the ERP has been accuratelymeasured. Relevant experimental parameters have also been tabulated.3.ERP UNDER REDOX POTENTIAL CONTROLClearly, when the potential of an aqueous solution of any redox couple followingprolonged irradiation is determined by hydrogen peroxide, the explanation for the factthat the ERP is identical with E(H202) and is independent of the redox couple and ofthe quality and intensity of the radiation is trivial. The remainder of the discussionwill therefore be confined to those cases where the ERP is redox controlled.In order to calculate the value of the ERP, it is first necessary to consider a modelsystem.(i) MODEL SYSTEM (UNIRRADIATED)Suppose that Hz02 is added to a dilute solution of a redox couple at a rate equalto GH202 x D' (eqn (2.10)) and that no He and OH- radicaIs are generated.One couldsuppose that the peroxide is run from a burette into a rapidly stirred solution. Thehydrogen peroxide will react according to reactions (2.8) and (2.9) and, under steadystate conditions,assuming that appreciable concentrations of MR+ and M("+l)+ are present and thatany unstable intermediates produced in (2.8) and (2.9) (e.g. OH., -H02)9 do not reactwith H202. The latter assumption is reasonable since OH- reacts preferentially withMn+, and the rate constant of the reactionis very small, lying within the range 0. I < k3.2 < 1 1. mol-l s-l at 25"C.lov2.8 = v2.9 (3.1)*H02 + H202+H,O+02 +OH* (3.274 EQUIVALENT REDOX POTENTIALThe net effect of eqn (2.8 )and (2.9) under steady state conditions is the catalyticdecomposition of hydrogen peroxide.The standard free energy change (= ph,o - pfiZo2) is thus independent of the activitiesof Mn+ and M(n+l)* i.e.of the redox potential (E (redox, model)). As indicated insection 2, the potential of inert electrodes immersed in aqueous hydrogen peroxidesolutions in the absence of other reacting species assume the hydrogen peroxidepotential at which reactions (- 2.2) and (2.3) occur at the same rates. The net effect isreaction (3.3) and the corresponding standard free energy change is also (phz0 -p;f202).If E(redox, model) # E(H202), part of the free energy of reaction (3.3) can be con-verted into electrical energy. Thus, part of the free energy of reactions (2.8) and (2.9)is converted to additional chemical energy during the establishment of the steady state.It is shown below that at least for the Fe2+/Fe3+ and 1--/12 systems,2H202 +2H20 + 0 2 .(3-3)E(redox, model) = E(H202). (3-4)(ii) KINETIC IMPLICATIONS OF EQN (3.4)It follows from eqn (3.4), (2.7) and (1.3) thatEo(redox) + (RT/F) In (ao/aR) = 0.835 + (RT/F') ln (aH+). (3.5)For eqn (3.5) to be compatible with eqn (2.1 1) and (2.12), the nature of the functionsf(0R9 a H + , aHzO2) and g(ao, aH+, must be such that= 1 under steady state conditions.Eqn (3.7) imposes some restraint on the choice of mechanism of reaction of H202 onthe components of the redox couple. The equation is verified for the Fe2+-Fe3+system in Appendix 11. However, it must be modified for more complex redoxcouples. Eqn (5.2) is the appropriate form for the Mn+ -M(n-tm)+ and eqn (A9,Appendix 11) for the 1--12 couple.From eqn (3.7) and (3.9,0.835 = (RT/F) In (k2.8/k2.9) + E, (redox).(3.8)It follows from eqn (2.14) of ref. (16) that if a, = aR,io(redox) = kzSl exp (-a FEo (redox)) = k-2.1 exp ((1 -a)FEo(redox)) i.e.where k2.1 and k-2.1 are proportional to the rate constants of the forward and reversecomponents of reaction (2.1). Eqn (3.8) and (3.9) imply that the product (k2.8/k2.9)(k2.1/k-2.1) is independent of the nature of the redox couple. The conditions underwhich this is true have been deduced from the version of the electron transfer theoryof R. A. Marcus *They are (1) that the rate determining steps of (2.8) and (2.9) involve a one electrontransfer to or from H20z; (2) that the reactions can be described by intersectingquadratic potential energy surfaces ; (3) that they are each characterised by a singleBrarnsted slope; and (4) that the sum of the Brarnsted slopes is approximately unity.Details are given in Appendix 111.Eo(redox) = (Wf3 In ( k 2 .1 k 2 . 1 ) (3.9)which is applicable to electron transfers with strong overlapP. L. AIREY 75(iii) IRRADIATED SYSTEMpH<3, eqn (1.1)) need to be considered in addition to those of H202.Under irradiation, reactions of the unstable intermediates (H- and OH- radicals at(3.10)OH-+Mn+-+M("fl)+ +OH-(+H++HzO). (3.11)If G H > GOH these reactions lead to net reduction in deaerated solutions. Understeady state conditions this is balanced by net oxidation by hydrogen peroxide ifappreciable concentrations of both redox components are present.A steady statecan be established if2 GH202) IGH-GOHI. (3.12)Supposefis the fraction of H202 reacting as an oxidising agent (eqn (2.8)) ; and(1 -f) the fraction reacting as a reducing agent (eqn (2.9)). Since hydrogen peroxideis a two electron oxidising or reducing agent it follows from eqn (2.8), (2.9), (3.10) and(3.1 1) thatthe net rate of oxidation = Go,D'+2GH,o,D'f, (3.13)and the net rate of reduction = GHD'+2GH20,D'(1 -f). (3.14)Under steady state conditions, the net rates of oxidation and reduction are equal.H. + M(n+ I)+ +Mn+Equating (3.13) and (3.14) and solving forf, one obtains,andThe velocity ratio of reduction and oxidation by H202 isf = (2GH202 + G H - GOH)/4GHzO2(1 -f) = ( ~ G H ~ o ~ - GH + G O H ) / ~ G H ~ O ~ .(3.1 5 )where C is a parameter which increases with increasing LET but is independent of theradiation dose rate at low solute concentrations and the nature of the redox couple.The linear energy transfer parameter, LET, is the energy deposited per unit lengthof the radiation track.aolaR = k2.8an+C/k2.9 (3.16)From eqn (3.15) and (3.6),where a.and a, are the activities of M@+ and Mnf . Under steady conditions,E = ERP = Eo(redox) + (RT/F) In (ao/aR) (2.7)= Eo(redox) + (RT/F) In (k2.8aH+ C/k2J. (3.17)From (1.3), (3.5) and (3.7)E(H202) = Eo(redox) + (RT/F) (k2.8aH + Ik2.9). (3.18)Comparing (3.17) and (3.18)(3.19)At O < p H < 2 , for low LET radiation G H = 3.65, GOH = 2.95 and GH2O2 = 0.8 l4 andtherefore, from eqn (3.15) C = 0.39.Substituting in eqn (3.17),ERP = E(H,O,) - 24 mV (0 <pH < 2). (3.20)For high LET irradiation:(e.g.GH = 0.25, GOH = 0.45 and G H ~ ~ ~ = 1.55.ERP = E(H202) + (RT/F) In C.'Po a-particles),76 EQUIVALENT REDOX POTENTIALThusC = 1.07 and ERP = E(H,Q2) + 1.8 mV (O<pH<2) (3.21)4. DISCUSSION OF THE PROPERTIES OF THE ERP I N TERMS OF THEABOVE THEORY(i) ERP IS NUMERICALLY CLOSE TO THE HYDROGEN PEROXIDE POTENTIALA N D IS ESSENTIALLY INDEPENDENT OF THE NATURE OF THE REDOXCOUPLEThe theory predicts that the ERP is given by eqn (3.19) provided (k2.1/k2.1) x(k2.8/k,.,) is independent of the nature of the redox couple (Appendix 111). TheERP is then independent of the nature of the redox couple since the parameter C(eqn (3.15)) is a function only of the primary radiation chemical yields.Followingirradiation, the potential will adjust to E(H202) provided there is sufficient hydrogenperoxide present to establish a new steady state. For 6oCo y-radiation, post irradia-tion effects of +24 mV are predicted (eqn (3.20)) in good agreement with the valueof +20 mV observed by Feates and Knight.6(ii) ERP I S INDEPENDENT OF THE INTENSITY AND THE QUALITY OF THEThe dependence of the ERP on the intensity and quality of the radiation will bereflected in changes in the value ofc = (2GH202 - GH f GOH)/(2GHz02 + GH - GOH). (3.15)Comparison of eqn (3.20) and (3.21) shows that the ERP is expected to increase byonly 25.8 mV in going from low LET to 210Po cc-particles.Variations in the radicalyields with dose rate are so small that they would not be reflected in the value of theERP. If the conditions were such that io(redox) was approximately equal to cO(H202),a dose rate dependent mixed potential between E(H202) and E(H202) + (RT/F) In C(eqn (3.19)) would be expected.RADIATION(iii) RADIATION CHEMISTRY OF REDOX DYESTUFFS CANNOT BE PREDICTEDDifficulties arise because the He and OH- radicals frequently react by addition andnot by oxidation and red~cti0n.l~ In many cases hydrogen peroxide does not reactaccording to eqn (2.8) and (2.9). The products, however, may still be electrochemi-cally active. The resulting steady state potential (if established) would probably bedetermined by the dyestuff, or its derivatives, but would not be related to the ERP.Complications would also arise due to the adsorption of the dyestuffs on the electrodesurface.FROM THE ERP5 .GENERALISATION OF THE THEORYIn the above treatment it is assumed that the radiolytic intermediates react asone electron reducing and oxidising agents respectively ; that HzOz is a two electronoxidising and reducing agent ; and that 2GHzO, > IGH - GoHl (eqn (3.12)).Clearly, the treatment is independent of the physical nature of the species (e.g. thereducing species could be e&). Moreover a fraction, 4, of the reducing radicals couldact as a one electron oxidising agent providedThe value of C (eqn (3.1 5)) and therefore the difference between the ERP and E(H,O,)would vary with 4.2GH,,,> 1(1 -#)GH-q5GH-GOHI(compare eqn (3.12)). (5.1P .L . AIREY 77The ERP will be independent of the presence of air despite the quantitativeoxidation of the Ha atoms, to *H02 radicals provided eqn (5.1) is obeyed whereFinally, the ERP will be established with redox couples of the type M"+, M(n+m)+G H = GHOz*provided the velocity ratio V2.9/V2.8 (eqn (3.6)) takes the form= 1 under steady state conditions.This is to ensure that substitution of (3.16) into the Nernst equationequation of the type (3.17).(5.2)leads to anThe author gratefully acknowledges contributions to this paper by Professor SirFrederick Dainton, F.R.S. and by Mr. D. F. Sangster and other colleagues at theA.A.E.C.APPENDIX I*THE NATURE OF THE POTENTIAL DETERMINING SPECIESFERROUS-FERRIC SYSTEMA potential of 0.84 V (= 0.87-0.059 pH) was observed on prolonged irradiation of aferrous-ferric solution (total concentration 5 x Theconcentration ratio [Fe3+]/[Fe2+] was found from the Nernst equation to be 300.Thus[Fe3+] = 5 X M and [Fe"] = 2x M at the ERP. The value of Eo (Fe2+- Fe")found experimentally in 0.4 M sulphuric acid (0.69 V6) was assumed. The exchange currentdensity, io (redox), was estimated to be A cm-2 using eqn (2.6) to extrapolate the resultsof Gerischer.18 The maximum value of the hydrogen peroxide concentration M) wascalculated, and the corresponding value of the hydrogen peroxide current io(H202) (2 x lo-'A cm-') was found experimentally. Thus io ( H 2 0 , ) Q io(redox) and the ferrous-ferric coupleis potential determining.M, pH 0.5, D = lo1' eV 1.-l s-').IODIDE-IODINE SYSTEMIt is possible to demonstrate by an analogous calculation that in acid solutions (pH<2)the ERP is determined by the redox couple.The calculation is complicated by the presenceof I; which undergoes electrode reactions at platinum and modifies the kinetics of the hydro-gen peroxide reactions in the bulk. By considering the rate at which the ERP is approachedduring irradiation from equilibrium potentials which are initially either below or above theERP, Henderson et aL4 were able to demonstrate that in lod4 M iodine solutions in 1 MH2S04, the ERP is redox controlled. The same, of course, would be true at higher iodide-iodine concentrations. In neutral and alkaline solutions, however, the nature of the poten-tial determining species cannot be decided, since there is no direct experimental evidence, andinsufficient kinetic data to permit calculation.OTHER SYSTEMSThe equivalent redox potential of a cerous-ceric solution (total concentration M,pH 0.37) is 0.79 V (against n.h.e.).4 The concentration ratio [Celv]/[CellK] is determinedfrom the Nernst equation (2.7) to be about i.e.[CeIV] 2: 10-l8 M and [CemJ = M.Because of the low [CeIV], io(redox) is very low (< 10-l2 A C I I - ~ ) . ' ~ Thus hydrogenperoxide is potential determining since its concentration builds up due to the virtual absenceof CeIV and its low reactivity towards Cem.* Following the practice in the references quoted, the distinction between concentration andactivity has not been made78 EQUIVALENT REDOX POTENTIALIt can be shown likewise, that, at the ERP the steady state concentrations of the oxidisedcomponents of the TIJ-Tllll couple and the reduced components of the 1-40; and thearsenious oxide system are less than 10-l2 M.Hydrogen peroxide is thus potential deter-mining in these cases.APPENDIX ITVERIFICATION OF EQN (3.6) FOR SYSTEMS EXHIBITING A REDOX DETERMINBDERPFERROUS-FERRIC SYSTEMThe discussion is based on the reaction scheme of Cahill and T a ~ b e . ~Fe3++H,02+FeOOH2++H+ (A0FeOOH2++Fe2++H02 ( A 3Fe3+ + H0,-,Fe2 + H+ + O2 (A3)Fe2++H202-+Fe3++OH+OH- (+H+-+H,O) (A41OH + Fe2 +-+ Fe3+ -k OH- ( + H++ H20) 645)Fe2++H02--+Fe3++HO; (+H++Hz02).(A61The process (Al) is assumed to be a rapid pre-equilibrium. The overall reaction (2.8)involves the steps (A4) and (A5) and reaction (2.9) involves the steps (Al), (M), (A3) and(A6). Reactions (A2) and (A4) are the rate determining steps. A fraction, #, of HOzradicals generated by reaction (A2) acts as a reducing agent (eqn (A3)), and a fraction (1 -#)as an oxidizing agent (eqn (A6)).The overall velocityv2.9 = kA2(1+#-(1-4)) [FeOOH2+] (A71= 2#k~,[FeooH+].Thuswhere KA1 is the equilibrium constant defined by process (Al). The fraction, 4, of radicalsreacting as a reducing agent is l/(l+k~~[Fe~+]/k~~[Fe~+]). It was seen in Appendix I that,at the ERP (pH 0.5), [Fe3++l/[Fez+] N, 300. The rate constant ratio is found by extrapolatingthe data of Barb et aLZ0 to be 5.The value of 4 is thus 1/(1+5/300) = 0.98, and, to a fistapproximation, is independent of the ratio [Fe3+J/[Fe2+] at potentials near the ERP.v2.9 = 24kA2KA 1 [Fe3+][H2021/aH -I-The velocityv2.8 = 2kA4[Fe2+][H2021*Thereforev2.9/ v2.8 = ( 4 k A 2 K.41 /kA4)[Fe3++]/[Fe2+’laH + (conlpare eqn (3*6)). (A8)IODINE-IODIDE SYSTEMLiebhafsky 21 has shown that the ratio of the rate of oxidation to reduction of Hz02in iodine-iodide solution is“21 x constant v2.9 - --V2.8 [I- I 2(% + l2Eqn (A9) follows directly from eqn (3) and (9a) of ref. (21). This is clearly the appro-priate form of eqn (3.6) for an electrochemical process of the typeI, + 2eg ~ 2 1 - (fwsince substitution in the Nernst equation (compare eqn (2.7)) leads to an equation of thetype (3.5)P.L. AIREY 79APPENDIX 111VERIFICATION THAT THE PRODUCT (k2.1/k-2.1)(k2.s/k2.9) IS INDEPENDENT OF THENATURE OF THE REDOX COUPLE FOR SYSTEMS WITH A REDOX DETERMINEDERPIt follows from eqn (3.9) that the heterogeneous rate constant ratio (l~~,~/k-~.~) is given byk2.Jk- 2.1 = exp (FE,(redox)/RT)= exp -(p&+-p;(n+1)+)/RT. (A1 1)The ratio (k2.8/k2.9) can be discussed in terms of the theory of Marcus " 9 l2 who showedthat the rates of homogeneous and heterogeneous electron transfer reactions can be expressedin the formwhere 2 is the collision number of the corresponding neutral species and AP* is the Gibbsfree energy barrier to the reaction measured from the energy of the reactant pair complex.22Thus from (All) and (A12),k = 2 exp (- AF*/RT) (A121(A13) k2.l k2.8 -k-2.1 k2.9-- - exp -(p&n+ -/.t&(n+j1+ +AF&-AF;.9)/RT.In its fundamental form the theory assumes weak overlap in the transition state.Thisclearly does not apply to the reactions of H202 which involve the making and breaking ofchemical bonds. However, the theory has been extended to atom transfer, proton transferand electron transfer reactions involving strong overlap. A functional relationship (eqn(21), ref. (13)) has been developed between the energy barrier AF*, the degree of reactionparameter n and the overall free energy of reaction AF". When applied to the electrontransfer reactionit takes the form AF214 = nAF~14+~(AFT,-oll)gl(n)+3(AF~2-co22) g2(1-n)whereAF*,,, and AFfi are the free energy barriers for reaction (A14) and the exchange Ox(i)+Red (i)+Ox (i)+ Red (i) ; cog is the energy of formation of the reactant pair complex fromreactants i, j at infinite separation; g(n) and g(1-n) are functions of n normalised so thatg&) = 1 ; and n+ is the value of n in the activated complex.The applicability of eqn (A15)which extends to reactions for which a quadratic potential energy function is appropriatehas been verified for a wide range of atom transfer and proton transfer reaction^.^^, 2 3 9 24The simplified hypothetical redox reactions (A16) and (A17) will be considered first.Complications which arise when a reactant in the rate determining step is generated in apre-equilibrium process involving H202 or another reactant are discussed when the treatmentis extended to the Fe2+/Fe3+ and 1-/12 systems.Ox( 1) + Red (2) -+ Ox(2) + Red (1)dAF&,/an = 0 (n = n")(A14)(A151MR+ + H202+M(n+1)+ + SR (A1 6 )M("'l)+ +H20z-+M"' +So (A171So and SR are the initial products of the one electron oxidation and reduction of H202.If(A16) and (A17) are the only rate determining steps of reactions (2.8) and (2.9),AFg.8 = AFz16 = n216AFi16 f~AFT1gl(n,f,,)+3AFT2gz(l -&6)AF;.9 = AFgl 7 = nz17AFi17 +~AF:lgl(n~17') + $AFS3g3(1 -nf17>.(A18)(A19)The subscripts 11,22 and 33 refer to the Mn+ /M("+ l)+, H202/SR and H202/S0 exchangesrespectively. The terms cog are normally much smaller than the free energy barriers and canbe neglected. nz is the value of the degree of reaction parameter for reaction ( x ) in theactivated complex.It lies between 0 and 1 and is close to 0.5 for weak overlap electro80 EQUIVALENT REDOX POTENTIALtransfer.13 The free energies of reaction A h F O are evaluated from the standard chemicalpotentials.A F i 1 6 = p& +p&+i)+ -pg202-p&n+ (A201A F i 1 7 = &,, +CLKn+-&,o,-PR(-+1,+From eqn (A13), (A18), (A19), (A20) and (A21),/AF;2g2(1 -nz16)-$AF:3g3(1 - nAfl ,))IRT (A22)provided + = 1. (A231This implies that deviations of n&, and n t l , from 0.5 are of the same magnitude but oppositesign. The value nxf is equal to the Brarnsted slope, 8AF*/8AF0.13 Thus ngI6 and nxl, areindependent of the nature of the redox couple only if the sets of reactions (2.8) and (2.9) ofwhich (A16) and (A17) are the rate determining steps are each characterised by a singleBrllrnsted slope.It follows from eqn (A22) that under these conditions (k2. /k-2.1)(k2,s/k2.9)is independent of the nature of the redox couple.I&- sYsmM.-Liebhafsky 21 has shown that the important rate determining stepsof reactions (2.8) and (2.9) are the electron transfer processesI-+ H202+IO-+ H2O (A241andIO-+H202+I-+H20 + 0 2 .The oxidised component 10- is different from, but in equilibrium with the oxidised compon-ent of the electrochemical couple, Iz.The modification of the treatment is quite general but will be applied directly to the I&system. The rate constant ratio k2.8/k2.9 takes the formI2 + H,O+I- + 10- + 2H+, KA2 6.(A261k2.81k2.9 = kA24/kA25Kk26 = (kA24/kA25) exp (2&+ $ - ~ ~ - + ~ & - - & - p H 2 0 ) (A27)where KA26 = K.426aHZ0 and p1-1~0 is the standard chemical potential of the solvent.in (Al3),When (A27) and the appropriate forms of (AlS), (A19), (A20) and (A21) are substituted(kz. 1 lk - 2,l)(k2.8/k2.9) = (k2. Ilk - 2. l)(kA24/kA2 5 K a Z 6)= exp - (nAz24(ccs", - &r2O2) - nA"2 dPi0 - P&OJ 4- + -pH20+*F%2g2(1 - nz24>-*F&g3(1 -. n,"z5))/RT* (A28)The required ratio is thus independent of the nature of the redox couple assuming, as above,that the Brarnsted slopes are equal i.e. that nXI6 = n&4 and n i l , = n i 2 5 .Fe2+/Fe3+ sYsTm.-It is widely agreed ' 9 2 5 9 that the rate determining steps areFe2+ + H202+Fe3+ + .OH + OH- (A29)andFe3+H0T -+(Fe2+H02)yFe2+ + HO2.Reaction (A29) is treated in the normal manner.During the course of reaction (A30), therewould be changes in the vibrational coordinates associated with the complex and its innersolvent layer as we11 as change in the orientation of the outer layers. The reaction couldtherefore be described by intersecting quadratic potential energy surfaces and functionalrelationships of the type (A1 5 ) would apply. Fe3fH0; is a precursor complex 22 and AF;2P. L. AIREY 81the barrier to the exchange reaction between species (e.g. HOT and H02), associated withHaOz. The free energy of the reaction is given byM i 3 0 = &k+HOt-&e3+HO;= (&e2+ +&,,+AF,"(Fe2+H02))-(pEe,+ +p& +AF,"(Fe3+HO;)), (A31)where AFf"(Fe2+H02) and AF;(Fe3+HO;) are the molar free energies of formation of thecomplexes from Fe2+ and HOz, and Fe3+ and HOT respectively.It follows from eqn (AS) and (A31) that(k2.dk2.9) = kA2914(= I K A i k ~ 3 0= (kA29/kA30) exp @;e3+HOi +&I' -P&0z-P;e3+)IRT= (k~zg/kA30) exp (AFi(Fe3'H0;)- AF;(H+HO,))/RT, (A32)where AFf(H+ HO;) is the energy of formation of H202 from H+ and HO; .By substituting(A31) and (A32) and the appropriate form of (AlS), (A19) and (A20) into (A13) it can beshown that( k2. 1 lk - 2. 1)(k2 .S/ k2 .9> = exp - (nlfi9(&~ - !'&oz) - nAf30(&bz - &$ -n~30AF~(Fe2+H0,) - nz29AFf"(Fe3 'HO;) + AF,"(H' HO,) +Thus, in those cases in which precursor (Fe3+ Hog) and successor (Fe2+ H02) complexes areformed, the rate constant ratio is only independent of the nature of the couple if the energiesof formation of the complexes are small compared with the Gibb's free energy barrier to thereaction.This is justified in the case of the Fe2f/Fe3+ system since the ERP has the valuepredicted by eqn (3.19). Insufficient information is available to relate eqn (A28) and (A33).All that can be said is that in both cases (k,.l/k--2.1#k2.sllC2.9) is independent of the natureof the redox couple under the conditions listed in section 3(iii).*Mg2g2(1 - nAfi9)-$AF:3g3(1 - n,f30))/RT. (A33)A. 0. Allen, J. Phys. Colloid Chem., 1948,52,479.F. S . Dainton and E. Collinson, Ann. Rev. Phys. Chem., 1951,2,99.V. I. Veselovskii, Int. Con$ Peaceful Uses of Atomic Energy, 1956,7, 599.I. H. S. Henderson, E. G. Lovering, R. L. Haines and E. J. Casey, Canad. J. Chem., 1959,37,164.G. H. Cartledge, Nature, 1960, 186, 370.F. S. Feates and B. Knight, Z h m . Faraday Soc., 1960,56,1680.J. O'M. Bockris and L. F. Oldfield, Trans. Fmuday SOC., 1955,51,249.A. E. Cahill and H. Taube, J. Amer. Chem. SOC., 1952,74,2312.lo F. S. Dainton and D. J. Currie, Trans. Faraduy SOC., 1965,61, 1156.l1 R. A. Marcus, Ann. Rev. Phys. Chem., 1964,15,155.l2 R. A. Marcus, J. Chem. Phys., 1965,44,679.l3 R. A. Marcus, J. Phys. Chem., 1968,72,891.l4 J. H. O'Donnell and D. F. Sangster, Principles of Radiation Chemistry (Edward Arnold,l5 M. Lefort and X. Tanago, J. Phys. Chem., 1959,63,833.l6 K. J. Vetter, Electrochemical Kinetics (Academic Press, New York, 1967), chap. 2, p. 117.l7 L. I. Grossweiner, Rud. Res. Reu., 1970,2, 345.l8 H. Gerischer, 2. Elektrochem., 1950,54, 366.l9 K. J. Vetter, Z.phys. Chem., 1951, I%, 360.2o W. G. Barb, J. H. Baxendale, P. George and K. R Ugrave, Trans. Furadzy Soc., 1951,47,462.21 H. A. Liebhafsky, J. Amer. Chem. SOC., 1932,54,1972.22 N. Sutin, Chern. in Brit., 1972,8, 148.23 A. 0. a h e n and R. A. Marcus, J. Phys. Chem., 1968,72,4249.24 R. A. Marcus, J. Amer. Chem SOC., 1969,91,7224.25 P. Jones, R. Kitching, M. L. Tobe and W. F. K. Wynne-Jones, Trans. Faraduy SOC., 1959,55,26 M. L. Haggett, P. Jones and W. F. K. Wyme-Jones, Disc. Faraday SOC., 1960,29,153.* R. Gerischer and H. Gerischer, 2. phys. Chem., 1956,6,178.London, 1970), chap. 6, p. 80.79

ISSN:0300-9599    

DOI:10.1039/F19736900070

出版商:RSC    

年代:1973

数据来源: RSC

摘要:

Electrochemical Behaviour of Mixtures of VO,’ and HzOz in1 M HC104 on Mercury and Platinized PlatinumBY L. NUCCI, R. GUIDELLIInstit Ute of Analytical Chemistry of Florence University, FlorenceAND G. RASPIInstitute of Analytical Chemistry and Electrochemistry of Pisa University, PisaReceirjed 3rd May, 1972The electrochemical behaviour of mixtures of VO; and H202 in perchloric acid media was studiedboth on platinized platinum and on mercury. From voltammetric measurements on platinizedplatinum it was possible to show that within a definite potential range (from 0.94 to 0.44 V/S.C.E.) theelectroreduction of VO(O2)+ to V02+ and HzO as well as its electro-oxidation to VOt and Oz arepreceded by the homogeneous decomposition of the VO(Oz)+ ion :PVO(Oz)+ + HZO+VO: + HzOz.The rate constant p = 0.74 s-I.The polarographic behaviour of VO(Oz)+ at a dropping mercury electrode in the presence of anexcess of H202 is characterized by the decomposition of this ion according to the preceding equation,followed by the electroreduction of VO: to VOZ+ and by the subsequent chemical regeneration ofVO(Oz)+ from V02+ by hydrogen peroxide.The regeneration reaction was shown to be first-orderwith respect to VO: and second-order with respect to HzOz ; its rate constant is 1.18 x lo3 1.2 moP.S-l.In preceding papers 1 s it was shown that hydrogen peroxide yields a compositewave characterized by a reduction limiting current of H202 to H20 and an oxidationlimiting current of H202 to O2 on platinized platinum (curve 1, fig. 1).Subsequentlyit was shown that the VO; ion in 1 M HCI04 yields a polarographically reversiblevoltammetric curve on platinized platinum ; this curve is due to the electroreductionprocessVO; +2H++e+V02++H20(curve 2, fig. 1). The aim of the present work is to study the voltammetric behaviourof mixtures of VO; and H202 in 1 M HC104, both on platinized platinum and at adropping mercury electrode.EXPERIMENTALThe mean polarographic and voltammetric currents were measured with a PolarecordMetrohm E 261 joined to an i.r. compensator model E 446. The current against time curvesat constant potential were determined with a potentiostatic system employing a Wenkingelectronic potentiostat model 66 TS 3. A Tektronix type 502 dual-beam oscilloscope was usedfor the measurement of instantaneous currents.A hemispherical platinum microelectrodeof radius ro = 1 mm with periodical renewal of the diffusion layer (DLPRE), described byCozzi and coworker^,^ was empIoyed in the voltammetric measurements. The platinummicroelectrode was platinized by cathodizing it for 4 s at a current density of 1.2 A cm-’ in a8L. NUCCI, R . GUIDELLI A N D G. RASP1 833 % solution of chloroplatinic acid containing 0.025 % of lead acetate. The droppingmercury electrode used in the polarographic measurements had the following characteristics :rn = 1.24 mg s-l, td = 5.75 s. All the measurements were performed at 25+0.1°C. Thepotentials reported in the following are referred to the saturated calomel electrode. Solu-tions were deaerated with purified nitrogen.VlV against S.C.E.FIG.1.-Voltammograms of 1.9 x M HzOz (curve l), 7.5 x M VO; (curve 2), and 2 xM VOi+ 1 x M HzOz (curves 3 and 4) in 1 M HC104 on platinized platinum.VOLTAMMETRIC BEHAVIOUR ON PLATINIZED PLATINUMRESULTSA solution of 1 M HC104 containing VO,* and H,02 in comparable amountsyields four consecutive voltammetric curves on platinized platinum. Their existenceis related to the formation of the monoperoxovanadate ion, VO(O2)+ :(1)Thus, starting from + 1.4 V and proceeding towards less positive potentials, in theneighbourhood of + 1 V we encounter the rising portion of an anodic irreversiblevoltammetric curve due to the direct oxidation of VO(O,)+ (curve 3, fig. 1) :VO(O2)++ H20+V02+ + O2 + 2H+ + 2e.The limiting current of this wave is controlled by diffusion.In fact, an oscillographicexamination of the corresponding instantaneous current shows that it obeys theCottrell equation corrected for sphericity :In eqn (2) t is the time measured from the start of electrolysis, D is the diffusioncoefficient of the reactant, and ro is the radius of the spherical electrode,VO(OJ+ + H,O+VO; + H20,.i = const (I /n% * + D*/ro). (284 ELECTROCHEMICAL BEHAVIOUR OF vo,' + H202As we proceed further towards less positive potentials we encounter a currentpeak due to the reduction of the platinum surface oxides, that obliterates the volt-ammetric curves furnished by the mixture of VO,+ and H202 at less positive potentials.In order to obtain these latter curves the platinum electrode was pre-cathodized at-0.6 V so as to eliminate surface oxides, then it was held for one minute at +0.1 Vin order to electro-oxidize adsorbed hydrogen and finally the voltammogram wasrecorded from +O.l V towards more positive potentials.At the four plateauxexhibited by the voltammograms so obtained (curve 4, fig. l), the following electrodereactions take place.(1) At the anodic plateau a both the hydrogen peroxide coming from the decom-position of the VO(O,)+ complex and any free H202 is electro-oxidized to 0,.The rate at which the decomposition of the VO(02)+ species proceeds partiallycontrols the anodic current at the plateau a. In fact, an oscillographic examina-tion of the corresponding instantaneous current shows that it does not obey(2) At the plateau b the VO,' ion, both free and deriving from the decompositionof V0(02)+, is electro-reduced to V02+.At the same time the free hydrogenperoxide, and that coming from the decomposition of VO(O2)+, is electro-oxidized to 0,. Obviously, the current at the plateau b also shows kineticproperties, on account of the relatively low rate at which the V0(O2)+ decom-position proceeds.(3) At the cathodic plateau c the VO,' ion, both free and deriving from the VO(O2)+decomposition, is still electro-reduced to V02+. At the same time the hydrogenperoxide, both free and deriving from the VO(02)+ decomposition, is electro-reduced to H20. The current at the plateau c also exhibits kinetic behaviour.(4) At the plateau d the peroxy-compound VO(0,) + is directly electro-reducedaccording to the overall electrode reaction :At the same time any free H202 or VO,' is also electro-reduced.In accord-ance with expectation, the current at the plateau d is diffusion-controlled, asapparent from the fact that the Cottrell equation (2) is satisfied.Of the three kinetic limiting currents corresponding to the plateaus, a, b and c,the one which is most suitable for studying the kinetics of the decomposition reactionof VO(O2)+ is the limiting current, El, at the plateau a, since along this plateau onlyH,02, of the three species VO(Oz)+, Hz02, and VO,', is electro-active. In orderto simplify the quantitative examination of the limiting current il it has been foundconvenient to render the bulk concentration, d*, of H202 negligible with respectto that, b*, of VO(O,)+.For this purpose, owing to the relatively high stability ofthe VO(02)+ complex, it is sufficient to keep the analytical concentration, CAY ofVO; slightly higher than that, C,, of H202 (say, CA32CD). In fact under theseconditions we may set b* = C' and d* = 0 to a good approximation. For values ofthe CA/CD ratio greater than 50, the limiting current il begins to become ill-defined.Consequently, it has not been possible to render the CA/CD ratio high enough tobe able to regard the VO,' concentration as practically uniform up to the electrodesurface during electrolysis.M.eqn (2)-VO(02)++4H++3e-+V02++2H,0.Table 1 reports i, as a function of CA for CD = 5 xDISCUSSIONThe gradual decrease of il with increasing CA, which is observed even for CA/CDratios greater than 50 (see table 1) permits us to establish that the dissociation reactioL.NUCCI, R. GUIDELLI AND G. RASP1 85of the VO(O2)+ complex (eqn (1)) is homogeneous. In fact, if this reaction tookplace in direct contact with the electrode surface, then, under limiting currentconditions, the rate of the backward reaction VO,' + H202-+V0(02)+ + H20should be vanishingly small on account of the zero value assumed by the surfaceconcentration of H202. This would imply the independence from C, of the netrate of reaction (l), and consequently of I,, at least when CA is high enough to causethe total complexation of the hydrogen peroxide present in the solution.TABLE 1 .-ii AS A FUNCTION OF CA FOR CD = 5 x M1 0 3 cA/moi I.- 1 il/,uA 103 CA/mol1.-1 ir/,uA1 .o 2.40 9.0 1.071.5 2.00 13.0 0.932.0 1.85 17.5 0.823.5 1.50 24.0 0.685.0 1.30 30.5 0.57Considering reaction (1) as homogeneous, the overall electrode reaction at theplateau a may be schematically depicted as follows :PPOB+A+ D (I), D+products+ 2e (11) (3)where B, A, and D denote the species VO(O,)+, VO,', and H20z whereas p and paare the rate constants of reaction (3-1) in the forward and bgckward directionsrespectively.The present diffusional problem was solved by Cizek, Koryta andKouteckf Subsequently,it was solved by one of the authors for diffusion towards a spherical stationaryelectrode, both with a rigorous method and with an approximate method based onthe diffusion layer concept.In view of the good agreement between the results ofthe approximate and of the rigorous method and since the approximate methodleads to an analytical expressin of il which is quite simple to use, in the presentwork we will employ this latter expression :for the case of diffusion towards a dropping electrode.z(id/i,+ = (il/id++)(2/d+ (4)In the preceding equation the dimensionless parameters z, 1, and 4' are defined asfollows :where a* is the bulk concentration of VO,', DD and DB are the diffusion coefficientsof H202 and VO(O2)+, f l is the period of electrolysis, and ro is the radius of thestationary spherical electrode. The parameter id = 2FA&b*[2/(&3$tf) + 1 /yo],where A is the electrode area, defines the anodic diffusion limiting current of VO(02)+.The diffusion coefficient DB of VO(O,)+ was derived from the experimental valueof the cathodic diffusion limiting current at the plateau d of curve 4 in fig.1 by usingthe Cottrell equation corrected for sphericity, and is equal to 5.6 x cm2 s-I.The diffusion coefficient DD of H202 equals 1.71 x cm2 s-I. Fig. 2 reportsexperimental values of (&/id + 11)(2/nit + 1/4°)2/(id + il - denoted by small circles,against DDtl/(DBb*). These values were obtained by changing the period of elect-rolysis tl while keeping botha* and b* constant (a* = 4.5 x M).Fig. 2 also shows experimental values of (illid + 1) (2/n* + 1 /#o)2 denoted by asterisksagainst, Dptl(id/i, - 1)2(DBb*).These latter values were obtained by changing the bulkM, b* = 5 86concentration a* of VOJ while keeping b* and t l constant (b* = 5 x M, t , =5.75 s). In agreement with eqn (4) the plot in fig. 2 is linear. From the experimentalvalue 1.85 x mol l.-I s-1 for the slope of this plot, which in view of eqn (4)coincides with p/a, and from the known value ' 4 x lo4 1. mol-I of the stabilityconstant a of VO(02)+, it is possible to attribute thevalue0.74 s-l to the rateconstant p .ELECTROCHEMICAL BEHAVIOUR OF VOt + H202-I-POLAROGRAPHIC BEHAVIOUR ON MERCURYRESULTSAs we have seen in the preceding section, the voltammetric behaviour of theVO; + H202 system on platinized platinum is characterized by the electro-activity ofH202 within the whole experimentally accessible potential range.On the contrary,Hz02 is electro-inactive on mercury at potentials more anodic than - 0.1 V. Further-more, the VO; ion is electro-reduced to V02+ on mercury at potentials more anodicthan the oxidation potential of mercury. Hence a wide potential range between+0.4 and -0.1 V exists within which VO; is electroreduced under limiting currentconditions, whereas H202 is electro-inactive.Kolthoff and Parry observed that solutions of Vv in 0.1 N H2S04 containing anexcess of H20z give a reduction limiting current from Vv to VIV which decreasesrapidly as the applied potential is shifted toward more negative values, so as togive rise to a current maximum. This limiting current is notably higher than thL .NUCCI, R. GUIDELLI AND G . RASP1 87diffusion limiting current of Vv to VIV as observed in the absence of H202, thusrevealing its catalytic nature. Kolthoff and Parry attribute this limiting currentto the electro-reduction of Vv to VIV, followed by the homogeneous regeneration of Vvby H202 according to a reaction of the type:Nevertheless, the above authors neither undertake a quantitative investigation ofsuch a limiting current (presumably on account of its remarkable dependence uponthe applied potential) nor do they establish whether the electro-active species is theVO(02)+ ion present in the bulk of the solution, or else the VO; ion coming from thedissociation of VO(O2)+.In perchloric acid ranging from 1 to 2 M the above catalytic mean limiting currentil varies only slightly with a change of the applied potential, provided the excess ofH202 is not too large (see fig.3). For sufficiently high concentrations of H202 the2 VXv + H202 4- 2H++2VV + 2H20.1-FIG. 3.-Polarogram of 2 x M H202 in 1 M HC104 at a dropping mercuryelectrode.limiting current of VO(O2)+ decreases appreciably with a shift of the applied potentialtowards more negative values. Confining ourselves to an examination of the polaro-grams for which il changes by no more than 10 % with a variation of potential from+0.35 to +0.25 V, we measured the dependence of il both upon the analyticalconcentration of Vv, CA, and upon that of H202, CD. The plot in fig. 4 shows theratio of the current il, as measured at +0.35 V, to the corresponding mean diffusionlimiting current, id, against C,.The data in fig. 4 were obtained from 2 x Msolutions of Vv in 2 M HC104 containing variable amounts of H202. For CD valuesless than M, the il/id ratio assumes values less than unity, exhibiting a minimumfor CD z 2~ As CD isincreased beyond this latter value, il, and consequently &/id, tend to increase linearlywith CD. Fig. 5 shows an experimental plot of it against CA for C, = 5 xThe instantaneous limiting current il at the chosen potential +0.35 V was measuredas a function of the time t elapsed from the beginning of the drop life, with an oscillo-scope. For values of the CD/CA ratio equal to or greater than unity the plot of log ilagainst log t is linear and exhibits a slope of about 2/3.M VO:+ 8 xM, namely for a unitary value of the CD/CA ratio.M.DISCUSSIONThe slope of the plot of log il against log t reveals that at the minimum of the illidagainst CD plot of fig.4 as well as along the subsequent rising portion the instantaneou88 ELECTROCHEMICAL BEHAVIOUR OF vo,’ -+ H202limiting current il is proportional to the instantaneous area A = 0.85 m* t* of themercury drop. It follows that the instantaneous current density at the droppingmercury electrode is independent of the electrolysis time t, thus revealing the rapidattainment of steady-state conditions. The preceding experimental observationspermit us to discard the hypothesis of a direct electro-reduction of VO(O2)+ onmercury. In fact, if this were the case, the initial decrease of iJid as CD is increasedcould only be attributed to a decrease of the diffusion coefficient of the electro-activespecies in passing from VO; to the corresponding coordination compound withH202, namely VO(O2)+.This interpretation must, however, be discarded in view ofthe fact that the diffusion coefficient of VO; , as determined from the diffusion limitingcurrent of this species in 1 M HCIOQ in the absence of H202 at a dropping mercuryelectrode, is equal to 6 x cm2 s-l and consequently is quite close to the diffusionlO~CD/rnOl I.-’FIG. 4.-Plot of Z& against lo2 CD for CA = 2x M.L. NUCCI, R. GUIDELLI AND G. RASP1 89coefficient of VO(O2)+. On the other hand, if the preceding interpretation werecorrect, the instantaneous limiting current il at the hydrogen peroxide concentrationcorresponding to the minimum of the i l / i d against C, plot should be diffusion-controlled and consequently proportional to t*, contrary to experiment.The kinetic nature of the limiting current at the minimum of the &/id against CDplot is to be ascribed to the low rate at which the electro-inactive species VO(O2)+generates the electro-reducible species VO:.The very presence of the minimum isattributable to the dual action of hydrogen peroxide which, on the one hand, bindsVO; yielding an electro-inactive species, and, on the other hand, regenerates VO;from the electrode product, V02+. The former action is felt mainly at the lowestH202 concentrations and leads to a decrease of &/id, whereas the latter action becomespredominant at the highest H202 concentrations and leads to an increase of illid.In the presence of an appreciable constant excess of H202, the mean limitingcurrent ir increases linearly with CA, as shown by the plot in fig.5. The precedingexperimental behaviour may be quantitatively interpreted by noting that the overallelectrode reaction under study proceeds through a mechanism of the type :PPUH20 + VO(OJ++VO; + H202VO; +2H++e-+V02++H202VO2++3H202+2V0(O2)++2H20+2H+ (III)( 5 )in which the order of reaction [(5)-(I)] is known, whereas that of reaction [(5)-(III)] isunknown. In the presence of a sufficient excess of H202, mechanism (5) can bedepicted schematically as follows :kd*mB $ A (I); A+e+C (10; C-+B (111) (6)where By A, and C denote the species VO(Oz)+, VO; and V02+.The rate constantpad* of reaction [(6)-(I)] proceeding toward the left embodies the concentration d* ofhydrogen peroxide, which is considered to be uniform throughout the solution up tothe electrode surface. The rate constant of the unidirectiona1 reaction [(6)-(11I)] alsoembodies the concentration d*, raised to the mth power, where m is the unknownorder of this reaction with respect to H202. Upon denoting by n the unknown orderof the same reaction with respect to C, it is possible to show (see Appendix) that, underparticular conditions, the expression of the limiting current il in the presence of asufficient excess of H202 is :pad*ii = +AoFt$ -Dkd*"Ci+l.Jn:1In the preceding equation td is the drop time, D is the diffusion coefficient common tothe species VO;, V02+, and VO(O2)+, and A . = 0.85 m3 is the area of the droppingelectrode at t = 1 s.Eqn (7) (which is derived in the Appendix) permits us to explain the experimentalbehaviour, provided we assume that reaction [(5)-(IU)] is first-order with respect toV02+ and second-order with respect to H202 (n = 1, m = 2) :il = &40Ft$JDkd*cA.In fact the experimental mean limiting current it is proportional to t$ as well as to theanalytical concentration, CA, of Vv.Furthermore, taking into account that th90 ELECTROCHEMICAL BEHAVIOUR OF vo,' -k H202cathodic mean diffusion limiting current of the VO; ion is given by the well-knownequationwe have :In agreement with eqn (8) the plot of i,/id against CD is linear when CD is much greaterthan CA and therefore CD E d*.Taking into account that in the measurements offig. 4 t d has been kept equal to 5.75 s, from the experimental value 67 1. mo1-l for theslope of this plot we deduce a value 1.18 x lo3 L2 mo1-2 s-' for the rate constant k.G. Raspi and L. Nucci, Ricerca Sci., 1967, 37, 509.R. Guidelli, L. Nucci and G. Raspi, Trans. Faraday Soc., 1968, 64, 3321.G Raspi and L. Nucci, Ricerca Sci., 1968, 38, 1054.D. CozZi, G. Raspi and L. Nucci, J. Electroanul. Chem., 1966, 12, 36.J. eizek, J. Koryta and J. KouteckL, Coll. Czech. Chem. Comm., 1959, 24, 663.R. Guidelli, J. Electroanal. Chein., 1971, 33, 291.M. Orhanovic and R.G. Wilkins, J. Amer. Chem. Soc., 1967, 89, 278.I. M. Kolthoff and E. P. Parry, J. Amer. Chem. Soc., 1951, 72, 5315.R. Guidelli, in Electroanalytical Chemistry, ed. A. Bard (Marcel Dekker, New York, 1971),VOI. 5, pp. 149-374.lo J. Kouteckjl and J. Koryta, Electrochim. Acta, 1961, 3, 318.APPENDIXThe diffusional problem relative to mechanism (6) may be solved by following the generalprocedure described in ref. (9), which is similar to that proposed by Koufeck$.lo If reaction[(6)-(1)] is sufficiently fast, then the amount of A which is consumed at the electrode surfacein a given time interval as a consequence of the charge-transfer step [(6)-(11)] is equal to theamount of A which is formed in the same time interval as a consequence of reaction [(6)-(I)] ina thin solution layer (reaction layer) whose thickness pl is much smaller than that of thediffusion layer.Under these steady-state conditions the differential equation satisfied by thespecies A in the neighbourhood of the electrode is given by :D(a2apx2> = pnd*a-pb for 0 < x < p1 (9)where a and b denote the concentrations of A and B, x is the distance from the drop surface,and D is the diffusion coefficient, which is regarded as equal for the species A, B, and C .Since pl is much smaller than the diffusion layer thickness, the concentration b may beconsidered to be uniform throughout the reaction layer (O<x<pl) to a good approximation.Let us denote by b the value assumed by b within this layer. Upon noting that (a/&)= 2(8~/8x)(~?~a/ax~), integrating eqn (9) between the limits x = 0 and x = pl yieldsthe relation :Under limiting current conditions we have a(x = 0) = 0.On the other hand, at x = p1reaction [(6)-(I)] is at equilibrium, so that the equality a(x = pl) = 6/nd* holds. Uponreplacing a(x = 0) and a(x = pl) in eqn (lo), we have L. NUCCI, R . GUIDELLI AND G . RASP1 91In deriving eqn (11) we have neglected (da/ax),, with respect to (aa/ax),=o, in view of thefact that the gradient of a must necessarily decrease in a gradual way in passing from x = 0Denotmg the orders of reaction [(6)-(III)] with respect to C and D by n and rn respectively,under steady-state conditions the concentration c of the species C in the neighbourhood ofthe electrode satisfies a differential equation analogous to eqn (9) :to x = pi:D(a2c/ax2) = kd*mc", for 0 < x < p2.(12)Eqn (12) holds within a thin solution layer bounded by the planes x = 0 and x = p2, wherep2 expresses the reaction layer thickness relative to reaction [(6)-(111)]. In practice, atx = p2 the species C is entirely consumed as a consequence of the oxidation of this species toB by hydrogen peroxide, so that we have :c(x = pJ r (ac/ax),, E 0. (13)Upon integrating eqn (12) between the limits x = 0 and x = p2 by a procedure analogous tothat followed in integrating eqn (9) and taking eqn (13) into account, we obtain :where co is the concentration of C at the electrode surface. At x = 0 the flux of A towardthe electrode, D(aa/ax),=o, is equal to the flux of C from the electrode, - D(~c/~?x),=o. Inview of eqn (1 1) and (14) we therefore have :On the other hand the instantaneous current is proportional to the flux of C from theelectrode according to the relation :il = - AoFt~D(ac/ax),=o.(1 6)Under steady-state conditions the concentration gradient (ac/ax),= 0 is time-independent, sothat the instantaneous current ii increases proportionally to t*. The mean current il isobtained by integrating eqn (16) between the limits t = 0 and t = td, where fd is the droptime :In order to obtain (ac/ax),=o, and consequently ii, as a function of the analytical con-centrations CA and CD z d* of Vv and H20z respectively, we must express co and b in eqn(15) in terms of these analytical concentrations. In this connection we note that, since thediffusion coefficients of the various diffusing species are practically equal, the followingrelation holds at any distance from the electrode :C, = a+b+c.(18)In practice two limiting cases may be encountered. In the fist case p1 @p2, whereas in thesecond one p2 &pl. If pl &p2, then from eqn (13) the concentration c is vanishingly smallalmost throughout the whole layer (O<x<pl) (see fig. 6a). On the other hand a is much lessthan b at any distance from the electrode. In fact, the equilibrium B+A is notably shiftedtoward the species B, particularly in the presence of an excess of hydrogen peroxide. Fur-thermore, in the neighbourhood of the electrode, A is consumed as a consequence of thecharge-transfer reaction [(6)-(II)], so that the value of the a/b ratio is less than that, (ad*)-l,Pp@d92predicted by the law of mass action.In view of eqn (18) we may therefore write to a firstapproximation :Replacing 6 from eqn (19) into eqn (15) and subsequently - D(dc/&),=~ from eqn (15) intoeqn (17) we have :According to eqn (20) il should decrease with increasing d*, in agreement with the fact that,under the present conditions, hydrogen peroxide is more active in binding the electro-reducible species A than in regenerating A from C . As a matter of fact the experimentalvalue of il increases with increasing C, E d* in the presence of an excess of H202, so that thecondition pl + p 2 is not satisfied.ELECTROCHEMICAL BEHAVIOUR OF vo,’ f H202-b z CA.(19)ir = +AoFt$CA(Dp/ad*)). (20)” I I-1 2FIG. 6.-Schematic representation of the concentration against distance profiles for p1 > p2 (a) andEL1 <P2 (b).In the case that pl<p2, the concentration c varies gradually from co to zero within alayer (O<x<p2) which is much thicker than the l_ayer (0< x < p l ) , within which the concentra-tion b remains practically uniform and equal to b (see fig. 6b). Consequently the concentra-tion of C may be regarded as constant and equal to co within the layer (O<x<pl). In viewof eqn (18) and taking into account that even in the present case a is everywhere negligiblewith respect to 6, we have :5 z c*-co. (21)Replacing 6 from eqn (21) into eqn (15) and rearranging, we have :Eqn (22) expresses co in an implicit form as a function of C, and d*. In order to obtain co inan explicit form, we shall limit ourselves to the consideration of two limiting subcases.In the first subcase, which is encountered when the equilibrium of eqn [(6)-(1)] is totallyshifted to the left, namely when o tends to co, we have L . NUCCI, R. GUIDELLI AND G . RASP1 93so that eqn (22) becomes :Replacing co from eqn (23) into eqn (14) and subsequently (ac/ax),=o from eqn (14) into (17)we obtain (20) once again. This is due to the fact that, provided Q is extremely high, themain effect of an increase in the concentration of HzOz is a corresponding increase in theentity of complex formation with production of the electro-inactive species B from theelectro-active species A, even if pl <p2. The remaining limiting subcase is expressed by theinequalityUnder the present conditions, from eqn (22) it immediately follows that co z CA. In viewof eqn (15) and (17) we therefore have :This is the only case, among those considered, which predicts the increase of il with increasingd* under the reasonable assumption that m is greater than zero. It should be noted that,while in the two other cases, the current ii depends exclusively upon the kinetics of reaction[(6)-(I)] by way of the parameter p , in the present case il depends exclusively upon the kineticsof the regeneration reaction [(6)-(III)] by way of the parameter k

ISSN:0300-9599    

DOI:10.1039/F19736900082

出版商:RSC    

年代:1973

数据来源: RSC

摘要:

Ionization of Moderately Strong Acids in Aqueous SolutionPart 2.-Further E.1m.f. Studies of the Dissociation of the Bisulphate Ion tBY A. K. COVINGTON,* J. V. DOBSON AND K. V. SRINIVASAN $Dept. of Physical Chemistry, School of Chemistry,Newcastle upon Tyne, Newcastle NEl 7RUReceived 13th June, 1972New measurements on the cell :Pt,H2 I Na2S04, NaHS04, NaCl I AgCl I Aghave failed to substantiate the original measurements made by Hamer forty years ago and sub-sequently reanalysed many times. The range of vaIues for the dissociation constant of the bisulphateion (K2 of sulphuric acid) is substantially in agreement with that obtained from the simpler cell inwhich sodium chloride was omitted and the mercury-mercury(I) sulphate reference electrode wasused. Kz = 0.0113+0.O005 mol kg-l at 25°C where the uncertainty arises from a range of reason-able choices for the ion-size parameter.Sulphuric acid and its second ionization constant (KJ, the dissociation constantof the bisulphate ion, have been the subject of a considerable number of investigationsbut with justification because of the importance of this electrolyte.The originale.m.f. studies of Hamer have been extensivelyreanalysed 2-5 but only the latter have been repeated.6 It is now well establishedthat the value of K2 obtained from e.m.f. studies is dependent on the ion-size para-meter in the Debye-Huckel equation used to estimate the activity coefficient terms.Covington, Dobson and Wynne-Jones studied a similar cell to that used by Hamer,hoping to achieve a simplification by using the mercury-mercury(1) sulphate referenceelectrode and thereby making the mixed electrolyte simpler by omission of chloride :For the same choice of ion-size parameter the results for K2 were higher than thoseobtained from Hamer’s cell :H2 1 NaHS04(m,), Na,SO,(m,), NaCl(m,) I AgCl I Ag.Reanalysis of the results on the acid-acid cell of Davies, Jones and Monk,2restudied by Nair and Nancollas Since theactivity term is different for each cell there is no reason to expect concordant valuesfor a given choice of ion-size parameter, nevertheless the spread is large (see table 2of ref.(4)). Particularly puzzling, however, was the large slope obtained in the re-analysis of Hamer’s data when the dissociation quotient, corrected for Debye-Huckel interactions (pKi) was plotted against ionic strength (see fig.2 of ref. (4)) incontrast to the smaller slopes from the other two cells. Further, Hamer’s resultst Part 1, A. K. Covington, J. G. Freeman and T. H. Lilley, J. Phys. Chem., 1970, 74, 3773.$ present address : University of Patna, Patna 5, India.94and of Davies, Jones and MonkH2 I NaHS04(m1), Na2S04(m2) I Hg2S04 1 Hg*H2 I H a HzS04 I Agcl I Aggave values falling between theseA. K . COVINGTON, J . V. DOBSON AND K . V. SRINIVASAN 95show curvature on such plots at ionic strengths above 0.2, which was unexpected.Accordingly it was decided to repeat Hamer's work.EXPERIMENTALThree series of measurements were made in which (i) ml = mZ = m3 (ii) ml = m3 # m2(iii) ml # mz # m3.In the first two series, buffer solutions were prepared from sodiumsulphate+ hydrochloric acid. In the third series anhydrous sodium sulphate and constantboiling distillates of sulphuric acid and of hydrochloric acid were used. The first series isan exact repeat of Hamer's conditions. Each solution was prepared separately and not bydilution of a stock solution.Sodium sulphate (A.R.) was recrystallised twice from conductivity water and dried at110°C. It was then heated to near fusion and allowed to cool in a desiccator over phos-phorus pentoxide. Its purity was checked by titration of the effluent from a cation exchangecolumn (H form) with standard sodium hydroxide in a nitrogen atmosphere using the Radio-meter automatic titration apparatus. Sodium chloride was recrystallised and purified in asimilar manner.The preparation of constant boiling acid distillates has been describedel~ewhere.~~ * The sulphuric acid composition was also checked titrimetrically.Hydrogen and silver-silver chloride electrodes were prepared as previously de~cribed.~,Bias potentials were checked before and after the experiments using comparison electrodesstored continuously in 0.01 mol kg- hydrochloric acid. Silver-silver chloride electrodesdeviating more than 0.02 mV from the standard electrode were discarded. Measurementsin the solution mentioned above lo versus hydrogen electrodes showed the standard potentialto be 222.37 mV at 25°C.Three-way cell vessels were employed with the compartments separated by a tap.Eachcell contained one hydrogen and two silver-silver chloride electrodes, whose bias was checkedin each solution. Cells were rinsed three times with solution before filling and placing in thethermostat bath controlled at 25.OOk 0.01"C. Nitrogen and hydrogen with adequatepresaturation were bubbled continuously through the appropriate cell compartments.Measurements were made with a vernier potentiometer (Cambridge Instrument Co. Ltd.)and sensitive galvanometer (Tinsley type 4500L). The standard cell was frequently checkedagainst an N.P.L. certificated laboratory standard. Pressure corrections were made in theusual manner.ll Cells once they had reached steady values, which usually took about anhour, maintained these (k0.03 mV) for two hours.RESULTSExperimental results are given in table 1.The e.m.f. ( E ) of the Hamer cell isgiven by :where k = (RTln 1O)/F, and y with appropriate subscripts represents activity co-efficients.mso4 = m2+mlj; mHso4 = m , - m H ; m,, = m 3 , mNa = rnl+2m2+ni3The individual ionic molalities are :and the ionic strength Z = 2mH + 3m2 + m, + m3.Introducing the equation :for each ionic species i, where p is termed the ion-size parameter, gives from eqn (1) :(E-E")/k = -log n~Hm3 +2A13/(1 +PI*). (3)The dissociation constant of the bisulphate ion is defined as :KL = mHj%04YHYS04 /tk?HS04]1HS0496 IONIZATION OF A C I b S IN AQUEOUS SOLUTIONTaking logarithms and substituting from (2) gives :m,(m, + m,) 4AI*P K ~ = -log +------- rn,-mm, 1+p1+ (4)where the prime on pK2 denotes the fact that since eqn (2) does not include a termlinear in I, pKi may be expected to vary linearly with I.An Algol programme waswritten for the English Electric KDF9 computer to solve iteratively eqn (3) for mHand I =f(mH) for various choices of the ion-size parameter p. For each p, valuesof pKi calculated from eqn (4), against I were fitted using a linear least squaresprogramme to yield pK2 values. The values k = 59.159 mV and A = 0.5107mol-3 kg* at 25°C were Values of pK; for p = 1 .O and 1.5 are given in table 1.Some values are plotted in fig. 1, where comparison is made with previous work.Results for pK2 from the various studies are collected in table 2.4.9998.5229.96317.96817.96729.80934.7005.5468.28213.91817.8302.0616.10510.15814.23422.2444.9938.5109.95817.95117.94729.77434.6625.6268.40214.11918.0871.8945.7949.70413.63321.3324.9998.5229.96317.96817.96729.80934.7002.3733.5445.9557.6282.0616.10510.15814.23422.24451 1.60488.69482.21457.75457.55436.72430.80528.8951 1.70489.97479.90550.88502.7248 1.01467.01448.49103mHPKi103mHPK5103mHPK51 0 3 ~ ~PK’,103mHPK5103mHPKi103mHPKi103mHPKi103mHPK’,103mHPKb103mHPK5103mHPKi103mHPKi103mHPKi103mHPK‘,1 0 3 ~ ~PKip = 1.03.6971.9195.7601.9126.5 181.91610.5591.90610.5981 A9716.1191.87718.0801.8833.9811.9255.5411.9188.4321.90810.2321.9101.7321.9294.3921.91 86.6981.9088.7901.90613.1321.890p = 1.53.6011.9435.5371.9356.2351.9389.8311.9269.9071.91714.6691.89416.3031.8973.8771.9465.3441.9387.9981.9279.6091.9261.71 11.9544.2621.9426.4071.9318.3061.92812.1541.91A . K.COVINGTON, J . V. DOBSON AND K. V. SRINIVASAN 97TABLE 2.-vALUES OF pKz FOR VARIOUS VALUES OF pp = 1.0 p = 1.3 p = 1.5 p = 1.7Hamer (recalc.) 1.983 1.996 2.004 2.01 1this work 1.929 1.946 1.954 1.964ref. (4) 1.936 1.952 1.963 1.972DISCUSSIONIt is clear from fig. 1 and table 2 that the new results for pK2 are considerablylower than those of Hamer and further that the slopes of the plots of pKk against Zare smaller than those obtained from reanalysis of Hamer’s results.The new resultsare more in accord with those of Covington, Dobson and Wynne-J~nes,~ althoughnot coincident for the same choice of p. However, concordent values can be ob-tained for slightly different choices of p , e.g. pK2 = 1.954 for p = 1.5 for this workand pK2 = 1.952 for p = 1.3 from the mercury(1) sulphate cell data (see also fig. 1).One might have expected the smaller p value for the solutions containing chlorideion in that the mean distance of closest approach should be smaller when chlorideions and the larger sulphate ions are mixed but this comment is naive and placestoo much faith in the physical significance of the ion-size parameter...- “.-..........-I .....................-\1.88 -\\ I I I I I \ f0.I 0.2Z/moI kg-lFIG. 1 .-Extrapolation of data from the Hamer cell : 0, this work ( p = 1.0) ; 0, this work ( p = 1.7) ; A, this work ( p = 2.5). (No attempt has been made to distinguish the points from the separateseries as these obviously fall on the same straight line.) - - -, results of Hamer ‘ ( p = 1.0) ; - - -,results of Covington, Dobson and Wynne-Jones ( p = 1.5). This figure is plotted on exactly thesame scale as fig. 2 of ref. (4) where experimental points are included for previous work.’.The new results show a scatter about straight lines of less than 0.1 mV, somewhatbetter than do Hamer’s results,l and it remains to examine reasons for the discrepancy.The new results were obtained from three different series of measurements, that isfrom three separate series of solutions with different molality ratios whereas Hamerused only one stock solution and diIuted it.His analysis of hydrochloric acid wasgravimetric for chloride and for sodium sulphate he weighed the residue afterevaporating to dryness and fusing the salt in a muffle furnace. Less satisfactory,however, was the technique of evacuating the solutions to expel dissolved gases whensomc loss of water could have taken place. We believe that an error in stock solutionmolalities is the most likely cause for the discrepancy. A similar explanation wasI 98 IONIZATION OF ACIDS IN AQUEOUS SOLUTIONsuggested by Prue and Read,12 who were unable to reproduce results for formic acidobtained l3 in the Sterling Chemistry Laboratory at Yale in the same period.Indeeddifficulty has been experienced in substantiating two other sets of data 14* l7 fromHarned's group.We conclude that pR2 = 1.945&0.015 or K2 = 0.0113+0.0005 mol kg-' fromthe present measurements where the uncertainty represents a range of p values of1.0 to 1.7. This does not preclude a somewhat higher value being considered moreappropriate, indeed for p = 2.5 the standard deviation of pK; against I plots has aclear minimum and pK2 = 1.995. In a following paper we shall present some newdata for K2 obtained from spectrophotometric studies.One of us (K. V. S.) thanks the University of Patna for study leave (1965-8).' W. J. Hamer, J, Amer. Chem.SOC., 1934,56, 860.C. W. Davies, H. W. Jones and C. B. Monk, Trans. Faraday SOC., 1952,48,921.W. J. Hamer in Structure of Electrolytic Solutions, ed, W. J. Hamer (Wiley, New York, 1959),p. 236.A. K. Covington, J. V. Dobson and W. F. K. Wynne-Jones, Trans. Faruday SOC., 1965,61,2057.H. S. Dunsmore and G. H. Nancollas, J. Phys. Chem., 1964,68, 1579.V . S. K. Nair and G. Nancollas, J. Chem. SOC., 1958,4144. ' C. W. Foulk and M . Hollingworth, J. Amer. Chem. SOC., 1923,45, 1220.K. Kunzler, Anal. Chem., 1953, 25,93.A. K. Covington and J. E. Prue, J. Chem. SOC., 1955,3696.lo R. G. Bates, E. A. Guggenheim, H. S. Harned, D. J. G. Ives, G. J. Janz, C. B. Monk, J. E.b e , R. A. Robinson, R. H. Stokes and W. F. K. Wynne-Jones, J. Chern. Phys., 1956,25,361;1957,26,222.I-). J. G. Ives and G. J. Janz, Reference Electrodes (Academic Press, 1961), pp. 95-6. '' J. E. Prue and A. J. Read, Trans. Faruday SOC., 1966, 62,1271.l3 H. S. Harned and N. D. Embree, J. Amer. Chem. SOC., 1934,56,1042.l4 H. S. Harned and W. J. Hamer, J. Amer. Chem. SOC., 1935,57,27.l6 W. H. Beck, K. P. Singh and W. F. K. WynneJones, Tram. Faraday SOC., 1959,55,331.l7 A. K. Covington, J. V. Dobson and W. F. K. Wynne-Jones, Truns. Furaday SOC., 1965,61,2050.W. J. Hamer, J. Amer. Chem. SOC., 1935,57,9

ISSN:0300-9599    

DOI:10.1039/F19736900094

出版商:RSC    

年代:1973

数据来源: RSC

摘要:

Catalysis by Polyelectrolytes of the Ag+, Hg2+, and T13+Induced Aquations of Co(NHJ5Br2+BY NORIO ISE AND YOSHINOBU MATSUDADept. of Polymer Chemistry, Kyoto University, Kyoto, JapanReceived 19th June, 1972The aquation reaction of Co(NH3)5Br2+ was studied kinetically with Ag+, HgZ+, and T13+ asinducing cations. Addition of polyethylene sulphonate, polystyrene sulphonate and polyphosphatemarkedly enhanced the reaction. The acceleration factor, ranging from 10 to lo4, was dependenton the inducing cations, polyelectrolytes, their concentrations and the concentrations of foreignsalts. The enthalpy of activation (AH*) and the entropy of activation (AS*) were decreased byaddition of polyelectrolytes in the case of Ag+ and Hg2+, as has been observed previously, whereasAH* was not influenced and AS* increased in the case of T13+.The behaviour of T13+ was ascribedto the strong hydration of this cation.Previously, we have shown that reactions between similarly charged ionicspecies are accelerated by oppositely charged polyions, and this effect is interpretableusing the Brarnsted relation in terms of the marked decrease of the activity coefficientof the activated complex in the strong electrostatic field of the polyions. Thisinterpretation was also applied to the polybase catalysis on an E2 reaction betweenchloromaleic acid or chlorofumaric acid and OH-.lC The validity of this interpreta-tion was clearly shown for a series of ionic reactions between oppositely charged ionicspecies such as the ammonium cyanate-urea conversion.I d In these investigations,relatively low polyelectrolyte concentrations were employed, to see whether the poly-electrolyte catalysis can be mainly accounted for in terms of the activity coefficientin combination with Brnrnsted theory, which is valid for systems deviating only slightlyfrom ideal behaviour. In this paper, we extend our measurements to much higherpolyelectrolyte concentrations, using Ag+ and TI3+, in addition to Hg2+ previouslystudied, as inducing agents for the aquation of Co(NH3),Br2+,Co(NH,),Br2+ + M"++ H20 + Co(NH3),H203+ + MBr("-l)+ (Mn+ : Ag+, Hg2+or TI3+)and discuss the influence of the valency of the inducing ions on the catalysis by poly-electrolytes.EXPERIMENTALMATERIALSThe complex [CO(NH~)~B~](NO,)~ is as used in a previous paper.'* The perchlorate[CO(NH&B~](C~O~)~ was prepared in a similar manner, using 70 % HC10, instead ofconcentrated HN03.(Found: H, 3.88; N, 16.05; C1, 16.10; Br, 18.14. Calc. for[CO(NH~)~B~](C~O~)~ : H, 3.58 ; N, 16.6; C1, 16.77 ; Br, 18.90 %.)The perchlorate salt had, in agreement with the literature, an value of 1 . 6 7 ~ lo4 at253 nm. Different anions in the cobalt complex had no influence on the rates under ourexperimental conditions. A solution containing Hg(C104)2 and HC104 was prepared bydissolving HgO (reagent grade) in an excess of HC104 (to suppress hydrolysis of Hg2+).9100 POLYELECTROLYTE CATALYSISThe acid dissociation constant for the Hg aquo ion pK, is 3.7. A Ag+ solution was preparedeither by dissolving AgN03 (Merck, reagent grade) in pure water or from newly precipitatedAg2O.A stock solution of T1(c104)3 was prepared by oxidizing TINOS (reagent grade)with bromine. T1203 was precipitated by treating the resulting solution with NaOH.The oxide was thoroughly washed and dissolved in HC104 solution. The Agf and Hg2+solutions were titrated with a 0.1 N NH4SCN solution with NH4Fe(S04)2 as internalindicator. The concentration of acid in the stock solutions was determined in the presenceof a large excess of NaBr by titration with NaOH with a mixed indicator of BromothimolBlue and Phenol Red. A known volume of the T P stock solution was diluted to givea solution in HCIO4 of < 0.5 M and added to a 2 % solution of KI. The liberated iodinewas titrated against thiosulphate.Dilution was necessary because concentrated solutionsof HC1O4 liberate iodine from a KI solution. To obtain a NaC104 solution, HC104 wasneutralized with NaHC03. Deionized water having a conductivity of 5 x 0-l cmdlor below was used in the present work. Sodium polyethylene sulphonate (NaPES) wasfrom the Hercules Powder Co., Wilmington, Del. The degree of polymerization was re-ported to be 770. Sodium polystyrene sulphonate (NaPSS) was given by the Dow ChemicalCO., Midland, Mich. The degree of polymerization was claimed to be 2500. Sodiumpolyphosphate (NaPP) was given by the Monsanto Co., St. Louis, Mo. The degree ofpolymerization was believed to be about 5000. Solutions of these polyelectrolytes werepurified by using cation- and anion-exchange resins.The polymer concentration wasdetermined by a potentiometric titration.KINETICSIn most cases, the reactions were followed by the stopped-flow method, using a rapidscan U.V. spectrophotometer (Hitachi Manufacturing Co., Tokyo, model RSP-2). TheU.V. photometer consisted of a Gibson type four-jet mixer and an observation cell with 1 cmoptical path length and a volume of 0.07 ml. The apparatus was used for reactions withhalf-times of 50ms or more. Equal volumes of a 10-4N HC1O4 solution containingCo(NH&Br2+ and a polyelectrolyte solution containing inducing cations were forcedthrough the mixing jet into the stop syringe. Just before the plunger hits a mechanical stop,it made contact with a trigger switch which actuates the horizontal time-basis sweep for theoscilloscope display.The display on the storage oscilloscope, which indicates the opticaldensity against time, usually started just before the reaction began and continued during theexpected period according to the preselected time-base setting. Slower reactions withhalf-times of 30min such as Ag+-induced aquation without polyions were followed by anordinary U.V. spectrophotometer (Hitachi Manufacturing Co., Tokyo, model EPS-3T).Although solutions containing polyions and reactant ions did not give precipitates for a fewhours after preparation under our experimental conditions, a slow increase of the opticaldensity was observed after the reaction was complete for the highest polyion concentrations.This might be due to an interaction between polyions and the tervalent products, namelyCO(NH~)~H~O~+, which introduced a small uncertainty (about 5 %) in the '' infinity "reading of the optical density (Dm).In such cases, the D , values at low concentration wereused for calculation. The disappearance of Co(NH3)5Br2f was followed by changes in theoptical density at 254 nm. Almost all rate measurements were carried out under pseudo-first-order conditions with an initial concentration of inducing reagents in tenfold excessover C O ( N H ~ ) ~ B ~ ~ + . The stopped-flow measurements were repeated three or four timesfor each set of solutions, the average deviation of the rate constant from the mean valuebeing usually f 5 %. The pseudo-first-order rate constant, k , was calculated from the slopeof the In (D- D)03 against t plot, where D is the optical density at time t.The second-orderrate constant, k2, was determined by k2 = k/[inducing ion]. The solutions in a water-jacketed cell were thermostated to +O.I"C.RESULTS AND DISCUSSIONThe value of k/ko (the ratio of the pseudo-first-order rate constants obtained inthe presence and absence of polyelectrolyte) is plotted against concentration (mol l.-INOR10 ISE AND YOSHINOBU MATSUDA 101of the polyelectrolyte in fig. 1-3. In all cases, the addition of a small amount ofpolyelectrolyte led to an enormous acceleration in the reaction rate. The followingpoints are specifically noted. First, generally speaking, the reaction rate increasedwith the polyion concentration when this concentration is low, passes through amaximum, and decreases on further addition of polyion.This behaviour is interpret-able by the Brarnsted equation in terms of the activity coefficients of the reactants andof the critical complex. According to previous measurements of the mean activitycoefficients of the component electrolytes in the ternary system H20 +sodiumpolyacrylate + NaC1,3 the mean activity coefficient of NaCl (y3) decreases at first atlow concentrations, passes through a minimum and then increases with increasingpolyelectrolyte concentration. If this concentration dependence is generally true,and should be more pronounced for higher valence electrolytes, it would give rise tothe above-mentioned concentration dependence of k/ko.(It should be noted thatsuch a concentration dependence was interpreted by Morawetz using the concepts of[NaPSS]FIG. 1 .-Dependence of the acceleration factor for the Ag+-induced aquation on PSS concentrationat 25°C. [Co(NH3),Br2+] = 5.5 x lo-, M, [Ag+] = 7.5 x M (kz0 = 0.0833 M-' s-I).I IrPOher1FIG. 2.-Dependence of the acceleration factor for the Hg*+-induced aquation on polyion con-centration at .25"C. [Co("H3),Br2+] = 6x lo-' M, [Hg"+] = 5 x M (& = 8.6, M-' s-l)102 POLY ELECTROLYTE CATALYSISeffective concentratio~i.~) Secondly, the three kinds of polyelectrolytes showeddifferent acceleration abilities. The order was PP > PES > PSS (see fig. 2). Thismay be ascribed to the difference in the spacial charge density of the polyions.Withcloser spacing of the charged groups, the activity coefficients of the reactants andcomplex would be lowered more. Thus PP which has the highest charge densitywas most effective and PSS least.In order to make a comparison of the inducing cations a set of rate measurementswas performed under the same conditions. Results are given in fig. 4. The orderof the acceleration factors is T13+ z Hg2+ 9 Ag+. If the polyelectrolyte catalysison ionic reactions could be interpreted only in term of the electrostatic interaction0 s -YFIG. 3.-Dependence of the acceleration factor for the TI3+-induced aquation on PES concentrationat 25°C. [Co(NH&Br*+] = 6x M, ITl3+] = M, [HC104] = 0.1 N (kZ0 = 1.63 M-I s-l).4.0 c1 1 1 - 5.0 -40 -30 -20log [PES]FIG.4,-Comparison of the polyion effect on the induced aquations at 20°C. [CO(NH~)~B~~+] =6 x M, [inducing cation] = 1 x equiv. L-l, [HC104] = 0.05 NNOR10 ISE AND YOSHINOBU MATSUDA 103between polyions and oppositely charged reagent ions, the T13+-induced aquationshould be more strongly accelerated than the Hg2+-induced one, as Gould observedfor electron-transfer reaction^.^ Since the acid dissociation constant of the TI aquoion pK, is 1.1: at least half of the TI ions exist in the form of Tl(H20)50H2+under our experimental conditions even in 0.1 N HCL04. It is reasonable, therefore,that on the T13+ system, polyions are as effective as on the Hg2+ system.The influence of inert salts was examined for the Hg2+ system, using NaC104.The result is shown in fig.5. The effect of acidity on the acceleration factor is givenin fig. 6. From these figures, it is clear that increases in NaC104 concentration andacidity result in a sharp decrease in the acceleration factor. Since it has beenknown that the Hg2+-induced aquation rate is independent of the pH of the solution4.0 -3.0-n s5 M 2.0-0 .-.11.0--3.0 -2.0 -1.0log INaC1041FIG. 5.-Dependence a the polyion catalysis on simple salt concentration at 30°C. [Co(NH3), "+I= 6.7 x M, [Hg2+] = 1.32 x M, [HC104] = 1.88 x M, [PSS] = 3.75 x equiv. l.-'.at pH below 3,2 the acidity increase in fig. 6 can be regarded as an increase in ionicstrength. Therefore, the results shown in fig. 5 and 6 may reflect the shielding effectby inert salt ions on the polyions.In this respect, it should be mentioned that theincrease in [ClOJ (fig. 5) complicates the situation because Hg2+ and ClO; form anassociated complex HgCIOz at higher concentrations. The dissociation constantof this ionic complex K is reported to be 1.0 x M.'The temperature dependence of the present reaction systems is interesting. Inprevious papers, we have pointed out that polyelectrolyte cataIysis in substitutionreactions between similarly charged ionic species is due to decreases in AH* and AS*(enthalpy and entropy of activation) whereas simple electrolyte catalysis (primarysalt effect) is due to increases in AH* and AS*. The rate enhancement in the poly-electrolyte case thus originates from the decrease in AH* and that in the simpleelectrolyte case from the increase in AS*.In the present reaction systems with Ag+and Hg2+, AH* and AS* decreased with addition of the polyelectrolytes as is shownin table 1. This is in accord with previous observations mentioned above. However,for the T13+-induced reactions, A H * was not influenced by polyelectrolyte additionwhereas AS* increased, as is shown in table 1104 POLYELECTROLYTE CATALYSIS[PSSIFIG. 6.-Influence of [HC1041 on the polyion catalysis at 30°C. [Co(NH,),Br2+] = 6.7 x M,[Hg2+l = 1.32 x M, [HC104] from top to bottom ; 1.88 x 9.6 x 9.6 x and1.88 x 10-1 N.TABLE 1 .-THERMODYNAMIC QUANTITIES FOR THE INDUCED AQUATION OF Co(NH&Br2+ WITHAND WWHOUT ADDED ELECTROLYTE *added conc.x 1041 A H * / AS+/ AG " 1electrolyte equiv. 1. - 1 kcal mol- cal deg- Irnol- 1 kcal mol-Ag77.5 x M)none 14.4 - 15NaPSS 0.300 12.3 - 193.00 3.3 - 4130.0 12.6 - 13NaN03 90.0 20.8 +8Hg2+(5x M in HC1O4, 0.01 N)none bnoneNaPES 0.1851.85NaPP 4.00NaPSS 4.00NaCIO, 20018.512.014.19.56.34.55.09.214.5- 16-7- 12- 18 - 23- 20-11-5T13+(lW3 M in HC104, 0.10 N)none 10.2 - 23NaPES 1.85 8.9 - 157.40 9.7 -1114.8 9.9 -918.5 9.9 -9none 11.4 - 1818.918.115.516.618.416.716.213.011.711.411.112.416.017.113.412.812.612.516.70 25°C ; [Co(NH3),Br2+] = 6 x M ; "0,) = 6.7 x N ; c [HC1O4] = 0.56 NNOR10 ISE A N D YOSHINOBU MATSUDA 105This “exceptional” behaviour of the T13+ systems invites comment.If wediscuss AS* in terms of solvation-desolvation of the activated complex, the AS*increase for the T13+ case may suggest that the polyelectrolyte addition results in(1) a stronger hydration of the initial state, or (2) a weaker hydration of the activatedcomplex than in the Hg2+ (or Ag+)-containing systems. In other words, reductionin the hydration number of T13+ caused by the polyelectrolyte addition is smaller thanthat for Hg2+ or A@ systems. We suggest that T13+ is more strongly hydrated thanthe other two cations. This interpretation is favoured by the half-life period (7) forthe exchange process of water molecules bound to the metal cations. Accordingto Eigeq8 the z value is s for Hg2+ at 25°C. No measure-ments have been reported for Ag+, but it would be reasonable to assume a value ofs for 2.The difference in the strengths of the interactions of water with T13+and Hg2+ (or Ag+) would be responsible for the liberation of less water moleculesfrom T13+ by polyelectrolyte addition than from Hg2+. This should cause the strongerhydration of the initial state for T13+-containing systems, as mentioned above. Inthis connection, it is interesting to note an experimental observation reported byPosey and Taube : in the case of the T13+-induced aquation an important source ofthe water molecule in the final product, [CO(NH~)~H~O]~+, is the hydration spheres for T13+ andof ~13+.9This work was supported by a Grant-in-Aid from the Ministry of Education.(a) N. Ise and F. Matsui, J. Amer. Chem. SOC., 1968,90,4242.(b) N. Ise, Nature, 1970,225, 66.(c) T. Ueda, S . Harada and N. Ise, Chem. Comm., 1971,99.(d) T. Okubo and N. Ise, Proc, Roy. SOC. A, 1972,327,413.(e) T. Okubo and N. Ise, in preparation.J. N. Brnrnsted and R. Livingston, J. Amer. Chem. SOC., 1927,49,435.T. Okubo, N. Ise and F. Matsui, J. Amer. Chem. SOC., 1967, SS, 3697.(a) H. Morawetz and B. Vogel, J. Amer. Chem. SOC., 1969,91,563;(6) H. Morawetz and G. Gordimer, J. Amer. Chem. Soc., 1970,92,7532;(c) J. R. Cho and H. Morawetz, J. Amer. Chem. SOC., 1972,!M, 375.E. S. Gould, J. Amer. Chem. SOC., 1970,92,6797.F. Bas010 and R. G. Pearson, Mechanism oflnorganic Reacrwnr, 2nd edn. (Wiley, New York,1967), chap. 3. ’ C. W. Davies, Progr. Reaction Kinetics, 1961,1, 161.M. Eigen, Pure Appl. Chem., 1963,6,105.F. A. Posey and H. Taube, J. Amer. Chem. SOC., 1957,79,255

ISSN:0300-9599    

DOI:10.1039/F19736900099

出版商:RSC    

年代:1973

数据来源: RSC