摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(4), 617-623 617 Electrochemical Study of the Heterogeneously Catalysed Reaction between N,N-Dimethyl-p-phenylenediamine and CO~~~(NH,),CI* at+ Monometallic and Bimetallic Surfaces of Silver and Gold Yao-Hong Chen and Ulrich Nickel" Institute for Physical and Theoretical Chemistry, University of Erlangen-Nurnberg, Egerlandstr. 3, D-91058 Erlangen, Germany Michael Spiro Department of Chemistry, Imperial College of Science, Technology and Medicine , South Kensington, London, UK SW72AY The mechanism of the heterogeneously catalysed reaction between N,N-dimethylp-phenylenediamine and Co"'(NH,)5C12+ at silver, gold and silver-on-gold (Ag/Au) discs has been studied by means of electrochemical methods. Both the mixed (or mixture) potentials and the mixture currents were determined by recording the current-potential curves of the reactants.Silver halide, formed during the reaction, was determined by sub- sequent galvanostatic reduction. The reaction at silver was strongly inhibited by the formation of silver halide whereas the reaction at gold was inhibited by adsorption of the organic compounds as well as by iodide. At Ag/Au discs almost no inhibition occurred. The explanation is that the reduction of the cobalt complex takes place predominantly at the silver surface whereas the simultaneous oxidation of the p-phenylenediamine occurs on the gold. In this way the formation of inhibiting silver halide is suppressed as well as the inhibition caused by the adsorption of p-phenylenediamine.In a heterogeneously catalysed redox reaction between an oxidant (Ox) and a reductant (Red), the metal acts as a con- ductor of electrons from Red to Ox. In contrast to normal heterogeneous reactions, no direct contact of the reactants at the surface of the metal need take Contrary to elec- trochemical reactions no external power supply is necessary to cause the transfer of electrons through the metal/solution interface. The general mechanism of reactions of this kind is Ieschematically displayed in Fig. 1. A suitable model system to study both the kinetics and the mechanism of such catalytic reactions is the oxidation of N,N-dimethyl-p-phenylenediamine (PPD) by Co"(NH3),C12+. This reaction has been thoroughly studied by taking as cata- lyst silver and gold discs as well as the colloids of these noble metals.'-, The reaction occurs irreversibly because of the rapid hydrolysis of the reduced cobalt ~omplex.~ The first Fig.1 Schematic representation of the mechanism by which a nobleoxidation product of PPD is p-semiquinonediimine (SQDI+), metal catalyses a redox reaction an intensely coloured ion radical. SQDI' may give p-quinonediimine (QDI+) either by further oxidation or by dis- chloride could be partially or completely reduced by p-prop~rtionation.~,~The most important redox reactions are phenylenediamines, a reaction which is used to develop pho- summarized in Scheme 1. tographic material. This process, too, depends on the kind Experiments carried out with both noble metal colloids and concentration of halide ions (X-) present in the solution.and large noble metal discs have shown that silver is a better Scheme 2 summarizes the most important reactions catalyst than gold,,** but at the surface of silver the reduction responsible for both the formation and the reduction of silver of Co"'(NH3),C12+ was found to become inhibited because of halide on the surface of silver. the oxidative formation of silver chl~ride.~~~*'~ Chloride rel- The higher the concentration of halide ions the more easily eased from the reduced cobalt complex according to eqn. (4) does silver halide formation OCCU~.~+~ Owing to the different was usually sufficient to slow down the reaction rate con- solubility products, bromide is more effective than chloride siderably.This silver halide formation corresponds to the but less than iodide. On the other hand, the reduction of photographic bleaching process. On the other hand, silver silver halide becomes more difficult with increasing concen- PPD + CO"'(NH,),C~~+ noble metal * SQDI' + Co"(NH,),Cl+ +SQDI' + Co"'(NH,),CI2 noble metal QDI' + Co"(NH,),Cl+ + H' PPD + QDI+ + H+ 2SQDI' c.--Co"(NH,),CI+ CO:~+ 5NH, + C1-Scheme 1 Ago + CO"'(NHJ,C~~~ Ag+ + Co1'(NH3),C1+ (5) Agf + X-*A@ (6) noble metal AgX + PPD b Ago + SQDI' + X-(7) Scheme 2 tration of halide ions. Iodide disturbs more than bromide and much more than chloride. The addition of halide ions there- fore strongly influences the silver-electrocatalysed reaction between N,N-dimet h yl-p- phenylenediamine and Co"*(NH,),C12'.' The gold-catalysed reaction should not be influenced by the addition of halide ions because no insoluble gold salts are formed.However, with iodide an inhibition due to the adsorption of iodide at gold was This effect has also been observed with related catalytic system^.'^^'^ Experiments carried out recently with a mixture of silver and gold colloids have shown a superadditive effect.8 Under certain conditions the inhibition could be almost completely suppressed. In order to obtain more information about the mechanism of reactions of this kind, the oxidation of N,N-dime th y1-p-phen ylenediamine with Co"'(NH ,), C12+ has now been investigated by using metal discs.This paper deals with electrochemical studies on both the chemical and the electro- chemical formation of silver halide at silver and Ag/Au discs. The formation of the silver salt was studied indirectly by sub- sequent electrochemical reduction.' *' The experimental results were compared with the data obtained from the current-potential curves of the reactants and discussed in terms of the theory of mixture potential and mixture ~urrent.~. Experimental Rotating silver and gold discs of 4 mm diameter were taken for the study of both the oxidative formation of silver halide and the recording of the current-potential curves of N,N-dimethyl-p-phenylenediamineand [Con1(NH3),C1]Cl2. The silver halide was formed either electrochemically or chemi-cally with the cobalt complex.The discs were carefully pol- ished for at least l min with 0.25 pm diamond paste before each experiment and then thoroughly washed with water. Usually a rotation speed of 800 rpm was chosen. The zero- current potential was measured against a calomel reference electrode (0.1 mol dm-3 KCl). The Luggin capillary was filled with 0.1 mol dm-, KNO,. The silver halide formed on the disc by chemical oxidation with [Co'n(NH,),Cl]Cl, was determined by subsequent galvanostatic red~ction,'.~ taking the disc as a cathode in an electrochemical cell with a plati- num counter-electrode and a calomel reference electrode (0.1 mol dm-, KC1). Usually 0.1 mol dm-, KNO, + 0.01 mol dmP3 KC1 was chosen as electrolyte in an electrochemical cell.The current-potential curves for both the oxidation of N,N-dimethyl-p-phenylenediamineand the reduction of Co"'(NH3),C12+ were recorded in the same electrochemical cell, the noble metal discs being taken as either anode or cathode, as appropriate. Ag/Au discs were prepared by inserting the gold disc as cathode into a solution of 0.001 mol dm-, AgNO, and applying a current of -100 FA for 0.5, 1, 2, 5 or 10 s. Thus, a total charge of 0.05 mC up to 1.00 mC was passed. A simple calculation shows that 0.10 mC deposits 6.2 x 1014 silver atoms which corresponds to about three silver monolayers on the disc of 4 mm diameter. The thickness of a monolayer is about 0.288 nm according to the diameter of silver." The electrochemical measurements were performed with the potentiostat-galvanostat HEKA PG 284.J. CHEM. SOC. FARADAY TRANS,, 1994,VOL. 90 The N,N-dimethyl-p-phenylenediamine2HC1, the halides and the buffer substances Na,HPO, and KH,P04 were p.a. re- agents from Merck. The cobalt complex [Co1n(NH,),C1]C12 was synthesized according to Schlessinger.'6 All solutions were prepared with doubly distilled water which had been degassed by ultrasound under water jet-pump vacuum and saturated with nitrogen. The phosphate buffers were prepared according to Ssrensen, but diluted 1 :2 with water in order to avoid the formation of silver phosphate. All experiments were carried out under an argon atmosphere at 20 "C. Results and Discussion Experiments with Silver and Gold Discs Fig.2 shows some current-potential curves for the electro- chemical oxidation of N,N-dimethyl-p-phenylenediamine.All curves of this kind were recorded at the gold disc in order to avoid any disturbance because of the formation of silver oxide. For comparison the measurements were also carried out at a glassy carbon electrode. Almost the same results were obtained. The position of the curves depended on the pH. With increasing pH the curves were shifted to more negative potentials until a limiting value was obtained which is near the first protonation constant of N,N-dimethyl-p- phenylenediamine (pK = 6.3)." The insert figure shows the pH dependence of the half-wave potential.'8 Fig. 2 also shows several current-potential curves for the electrochemical reduction of [Co1''(NH,),C1]C12 .The results recorded at the silver disc electrode are indicated by 'S', while those obtained at the gold disc electrode are designated with 'G'.The position of these curves does not depend on the pH but does depend strongly on the kind of electrode material and on the kind and concentration of added halide ions. The higher the concentration of a given halide the more the curves shift cathodically. It is obvious that with silver as electrode the formation of silver halide is responsible for this shift. The electrochemical reactions for both the reversible oxida- tion of PPD and the reduction of Com(NH,),C12+ are sum- marized in Scheme 3.PPDI SQDI' + e-(84 SQDI' QDI' + e-(86) Co1*'(NH3),CIZ'+ e--Co"(NH,),Cl+ (9) Scheme 3 kc -40t --400 -200 0 200 400 ElmVvs. 0.1 rnol dm-3 calomel electrode Fig. 2 Current-potential curves of both N,N-dimethyl-p-phenylene-diamine (5 x rnol drnT3) at three different pHs and CO"'(NH,),CI*+ (5 x rnol dmP3) at pH 6.0 in the presence of the additional halide on a rotating silver disc (Sl-S4)or a rotating gold disc (G144);800 rpm, scan rate 10 mV s-'; (-) no addi-tional halide; (-.-.) 0.1 rnol dm-3 KCI; (---) loP3 rnol dm-3 KBr; (-+ .) lop3 rnol dm-3 KI. (a)pH 4.5, (b) pH 5.6, (c)pH 7.0. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Addition of eqn. (8a)and (9) gives eqn. (1) in Scheme 1, while addition of eqn. (8b)and (9) gives eqn.(2).The reactions (8) and (9) also take place at the surface of the noble metal in the absence of an external power supply if both redox couples are present in the solution together. The noble metal surface then takes up a mixed or mixture potential (Emix)so that it can act simultaneously as an anode for couple (8) and as a cathode for couple (9). An example is shown in Fig. 2. The mixture potential should not be confused with the zero-current poten- tial which is measured for a single redox system. At Emixthe net current is zero, owing to a balance between the anodic current resulting from the oxidation of PPD and the cathodic current produced by the reduction of Co"'. These two currents, termed mixture currents (Imix),are indi- cated by dashed vertical lines in Fig.2. By Faraday's law, Imix is directly proportional to the reaction rate at the surface (vc,Jmol m-2 s-') and is given by eqn. (I):2*14~1g ucat = Imix/nFA (1) where F is the Faraday constant, n the number of electrons and A the surface area of the catalyst (in m2). As long as the initial concentration of PPD exceeds that of the cobalt(r1r) complex the mixture current is likely to be in the region of the limiting current of CO"'(NH,)~CI~+. Under these condi- tions the kinetics were consistent with diffusion control, and both the mixture current and the measured reaction rate were proportional to the [CO"'(NH,),C~]~ concentration but + independent of the concentration of N,N-dimethyl-p-phenyle- nediamine.The heterogeneous reaction rate at both silver and gold discs is then given by eqn. (11):2*5 uCat= khct[Co"'(NH3)5C12t] (11) The curves displayed in Fig. 2 clearly show that the mixture current, and thus the heterogeneous reaction rate constant (k,,Jm s-'), depends on several factors. Most important is the kind of noble metal used as catalyst but also the kind and concentration of halide can strongly influence the position of the mixture potential and the value of the mixture cur- rent. Some values of the mixture potential and the mixture current, calculated from current-potential curves, are listed in Table 1. For the system presented in Fig. 2, the maximum mixture current and therefore the maximum reaction rate were obtained with the gold disc at pH > 4.5 (as can be seen by comparing the lengths of the dashed vertical lines).At the silver disc this reaction rate was obtained only at pH 2 7 and in the absence of additional halide ions. The initial reaction rate, however, was high also at a lower pH. The decrease of the reaction rate, i.e. the amount of inhibition, was found to depend on the relative position of the current-potential curves of PPD, Co"'(NH ,),C12+ and silver halide. The shift of the current-potential curves Sl-S4 to more negative potentials is caused by the formation of silver halide. Inspection of the curve S1 in Fig. 2 shows that in the absence of additional halide ions the formation of silver chloride began at about + 100 mV (us. 0.1 mol dmP3 calomel electrode).The necessary chloride ions were provided by the counter-ions of the salts employed. As at pH > 7 the mixture potential of the p-phenylenediamine and cobalt complex redox couples was considerably more negative, almost no inhibition occurred, but in the presence of additional halide the shift of the current-potential curve of Co"'(NH,),C12 to+ more negative potentials caused a dramatic decrease of the mixture current (cf:Table 1). The more negative the resulting mixture potential, the more difficult was the reduction of silver halide by PPD, because the mixture current for this catalytic process also became very small. In the presence of iodide the mixture current becomes zero. Indeed, the small amount of oxidized p-phenylenediamine present in the solu- tion was then able to oxidize silver! Particularly strong inhi- bition of the silver-catalysed reaction between PPD and Co1"(NH,),C12+ took place at low pH (because of the posi- tive shift of the current-potential curve of the p-phenylene- diamine), in the presence of a high concentration of chloride (shifting the current-potential curve of the cobalt complex to more negative potentials due to the formation of silver halide), and even more so in the presence of bromide and iodide (which shifted the cathodic current-potential curve to still more negative positions).The influence of additional halide on the current-potential curves of CO*~'(NH,),C~~+ recorded at a gold disc electrode is also displayed in Fig.2 (curves G144). In contrast to silver, the mixture current was not decreased by the addition of chloride and bromide, because Emixremained in the region of the limiting current for Co"'(NH3),ClZ+ except when the pH became lower than 4 (see also Table 1). However, as with the silver curve S4, the mixture current decreased strongly in the presence of iodide (curve G4). This was probably caused by inhibition of the electron transfer by adsorption of iodide on the surface of g0ld."*'~9~~ The desorption of iodide only occurred at a rather negative potential. However, not only iodide but also N,N-dimethyl-p-pheny- lenediamine itself and its oxidation products inhibited the catalysis due to an adsorption effect." The current-potential curves G2 and G3 in Fig.3 show the strong shift of curve G1 to more negative positions after a short treatment of the gold disc with the solution of the p-phenylenediamine. The corresponding values of Emixand Zmix obtained from these curves are summarized in Table 2 (left-hand column). The curves SG refer to Ag/Au experiments and will be discussed later. Because of the decrease of the mixture current the rate of the catalytic reaction between N,N-dimethyl-p-phenylene-diamine and Co1"(NH,),C12 should also decrease. In fact, + this decrease was found both at large gold discs and with colloids.2,8 In the absence of additional halide and at pH > 5 the reaction rate between PPD and Co"'(NH3),C12+ was usually considerably slower on the surface of gold than on the surface of silver.2*8 Table 1 Values of Emirand Imix under different pH conditions and with additional halide ([Co"'], = 0.5 mmol dm-3 and [PPD], = 0.5 mmol dm-3 at 800 rpm) Ag disc Au disc PH [KCl]/mol dm -[KBr]/mmol dm -[KI]/mmol dm -Emix/mV 'mixlPA EmixImV ImiJPA 4.5 --95 4 210 28 5.6 --90 15 140 28 7.O --65 28 65 28 6.0 --90 25 90 25 6.0 0.1 --20 5 90 25 6.0 ---50 4 90 25 6.O -1.o none 0 -85 2 J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 I 0.3f : 'I ................. Q: 1 -20 ...... . i-401 -1 I 1 1 -400 -200 0 200 400 E/mV vs.0.1 rnol dm-3 calomel electrode Fig. 3 Influence of the adsorption of N,N-dimethyl-p-phenylenedia-mine on the mixture current of the redox system N,N-dimethyl-p- phenylenediamine (5 x rnol dm-3)-Co111(NH,),C12+ (5 x mol drn-,) at pH 6.0 on a rotating Au disc (G143) and a rotating Ag/Au disc (SG1-SG3); 800 rpm, scan rate 10 mV s-'.The Ag/Au disc was obtained by inserting a gold disc into a solution of lo-, mol dm-, AgNO, and galvanostatically reducing it for 10 s at -100 PA. Both Au and Ag/Au discs were pretreated with 0.1 mol dm-, PPD for 0 s (full lines, G1 and SGl), 10 s (dashed lines, G2 and SG2) and 30 s (dotted lines G3 and SG3), respectively. Experiments with Ag/Au Discs Fig. 4 shows some current-potential curves for [CO"~(NH,),C~]C~, recorded at several Ag/Au disc elec- trodes. At very negative potentials only the reduction of the cobalt complex took place, but when the potential became more positive, a parallel reaction set in: the oxidative forma- tion of silver chloride.Thus, the cathodic current decreased -200 0 200 400 E/mV vs. 0.1 rnol dm-3 calomel electrode Fig. 4 Influence of both the amount of deposited silver and the addition of chloride ion on the electrochemical reduction of Co1"(NH3),C12+ (5 x lod4 rnol drn-,) at pH 5.6 on Ag/Au discs rotating at 800 rpm, scan rate 10 mV s-'. The amounts of silver deposited on gold discs 1-3 were 0.05,O.lO and 0.20 mC, respectively. (-) No additional halide; (---) 0.1 mol dm-, KC1 (with 0.10 mC silver on the Ag/Au disc). Insert: Dependence of the AgCl formed (QAICI)on the amount of silver (QAI)obtained from curves 1-3. QAga was calculated from the total areas of the corresponding peaks.thin silver chloride layer, the electrocatalytic reaction +between PPD and Co"'(NH3),C12 can therefore remain diffusion-controlled. The limiting current which under certain conditions (e.g. at pH > 5 and with an excess of PPD) equals the mixture current (see Fig. 2) is given by the relation (111) llim nFD, A[CO"'(NHJ,C~~']/a (111)= where D,, (in m2 s-') is the tracer diffusion coefficient of the cobalt complex in the phosphate buffer and or the net current even became anodic. Because of the limited is the concentration of the cobalt complex '3 [Co"'(NH3),C12 amount of silver deposited on the gold disc this oxidative formation of silver chloride stopped when all the silver had been transformed into silver chloride. As the potential rose still further, the only electrode process was again the reduction of the cobalt complex.The more silver had been deposited on the gold disc electrode, the higher was the resulting anodic peak. The amount of AgCl could be deter- mined from the total area of each peak in these curves; it was proportional to the amount of silver deposited on gold (see insert in Fig. 4), but independent of the concentration of chlo- ride present in the solution or released from the cobalt complex. This is shown by the dashed curve, which corre- sponds to the conditions of Fig. 4, curve 2, but was recorded in the presence of 0.1 mol dm-3 KC1. The addition of chlo- ride only shifted the peak cathodically, i.e. the formation of silver chloride began at a lower potential.As the formed silver chloride is porous,1othe deposition of AgCl on the gold disc decreased the limiting current (Imin)for the reduction of Co"'(NH,),Cl2+ only a little. In spite of a Table 2 Values of Emix and Imixobtained from the current-potential curves displayed in Fig. 3 none 90 25 90 25 10 0 6 90 25 30 -40 3 50 20 Discs were pretreated with 0.1 rnol dm -N,N-dimethyl-p-phenyle-nediamine (PPD) at pH 6. t is the pretreatment time of the disc with PPD. in the bulk. At a rotated disc the diffusion layer thickness 6 (in m) is given by the Levich equation:22 6 = 0.643D~.3v'/6f-1/2 (IV) where v is the kinematic viscosity of the solution (in aqueous solution ca.m2 s-') and f is the rotation frequency of the disc (in s-'). Fig. 5 shows some examples for the dependence of the lim- iting current (Ilim)on the square root of the rotation fre- quency (J) for the reduction of Co(NH3),C12+ at an Ag/Au disc. The value of the diffusion coefficient (Dco)was obtained 40 20L4 t I 1 -200 0 200 400 €/mV ws. 0.1 mol dm-3 calomel electrode Fig. 5 Dependence of the limiting current (Ilim)at two potentials on the speed of rotation of a rotating Ag/Au disc for the reduction of Co"'(NH3),C12+ (5 x rnol dm-3) at pH 5.6; scan rate 10 mV s-'.The amount of silver deposited on the gold disc was 0.10 mC. (a) 400, (b) 800 and (c) 2000 rpm. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 from eqn. (HI), taking Zlim at -100 and and +220 mV, respectively (indicated by the dashed line).At -100 mV a value of D,, = 7.16 x lo-'' m2 s-' was obtained, and at +220 mV the slightly lower value of 6.63 x lo-'' m2 s-'. The first value corresponds very well to the values obtained in previous papers in the absence of any phosphate buffer (7.24 x lo-'' m2 s-l lo and 7.30 x lo-'' m2 s-' '). As the diffusion coefficient should not depend on the potential of the electrode in this potential range, it can be assumed that the electron transfer through the interface of Ag/Au and of AgCl/Au is similar to that through the surface of pure silver or gold2,' and that the deviation in the value of D,, is only caused by a slight decrease of the active surface area of the disc.Similar results were obtained when the concentration of [Co"'(NH,),Cl]Cl, was varied between 0.2 and 0.5 mmol drn-,. The amount of electrochemically formed AgCl did not depend on the concentration of [Co"'(NH,),CI]Cl, ,but only on the amount of silver deposited on the surface of gold. In accordance with eqn. (111), the limiting current was pro-portional to the concentration of the cobalt complex. By combining eqn. (I) and (11), the rate constant k,,,/m s-could be calculated: khet = ZmiJ( n FA[Co(NH,) $1 I} (V)+ A value of khet = 5.0 x lo-' m s-' was obtained. This value is somewhat higher than that of 3.6 x lo-, m s-l reported previously.2 The difference can be explained by the slightly higher value of the mixture current obtained at the Ag/Au disc, because it more closely approached the limiting current for the reduction of the cobalt complex. It should be men- tioned, however, that this value is obtained only if the mixture current equals the limiting current of the cobalt complex reduction, i.e.in the presence of an excess of PPD. Fig. 6 shows some current-potential curves for the t i 11 3 -20-5 0 \ -40 --80 0 0.2 0.4 0.6 0 Q,,/mC -400 -200 0 200 400 €/mV vs. 0.1 rnol drn-3 calomel electrode Fig. 6 Influence of the amount of silver on the electrochemical reduction of silver chloride on a rotating Ag/Au disc in a solution of 0.1 rnol dm-, KNO, + 0.01 mol dm-3 KCl; 800 rpm, scan rate 10 mV s-'. Current-potential curves of the reduction of silver chloride started from the zero-current potentials (ca.300 mV) which were measured at the beginning. The amounts of silver deposited on the gold discs were 0.06 mC (curve l), 0.12 mC (curve 2), 0.21 mC (curve 3) and 0.60 mC (curve 7), respectively. (a) pH 4.5, (b)pH 7. Silver chloride was formed by the chemical oxidation of the silver with Co"'(NH,),Cl, (5 x mol dm-,) at pH 6.0 for 15 s. Insert: Dependence of both the zero-current potential (E")and the amount of AgCl formed (QA,,ci) on the amount of silver (QA,) on an Ag/Au disc. QADCl was calculated from the total area of each peak in the figure. 621 reduction of silver chloride deposited on an Ag/Au disc and for comparison the current-potential curves for the oxidation of N,N-dimethyl-p-phenylenediamineat two different pH values.The silver had been deposited on the gold disc gal- vanostatically as described above and the silver chloride was formed by subsequent chemical oxidation of the silver with Co"'(NH,),ClZ +. The more silver had been deposited, the more silver chloride could be formed during the chemical oxi- dation, but as the oxidation time was kept constant at 15 s, any large amount of silver deposited on the gold disc could not be completely converted into silver chloride. In order to avoid any dissolution of the silver chloride formed, the Au/ Ag/AgCl disc was then rinsed with water containing lop3 mol dm -potassium chloride. The current-potential curves were carried out in a cathodic direction, starting from the zero-current potential.It can be seen from the insert figure that the zero-current potential depended on the amount of silver deposited on the surface of gold. The silver chloride formed on the Ag/Au discs was deter- mined from the total area of the cathodic peaks in Fig. 6. When only a small amount of silver had been deposited on the surface of the gold, the silver was almost completely con- verted into silver chloride by the subsequent chemical oxida- tion. Under this condition, the activity of the residual silver (ffAg) was much smaller than unity because of the dominant gold surface. If, however, a relatively large amount of silver had been deposited, so that the silver had not been com- pletely converted into silver chloride, the activity of the silver could approach the value for pure silver, i.e.ffAg % 1. Assuming the validity of the Nernst equation for this gold/ silver/silver chloride electrode, the zero-current potential (EAg/Ag+) is given by RT RT--In aAgo--In ax-(VIb)F F where R is the gas constant, T the temperature in K and Leg, the solubility product of AgX. The zero-current potential, EAgO,Ag+, at the Ag/Au disc can therefore be much more posi- tive than at the surface of pure silver if the amount of silver is very small. In fact, the insert in Fig. 6 shows a strong depen- dence of the zero-current potential on the amount of silver deposited on gold. The points 1-4 correspond to the situ- ation where the chemical oxidation time was sufficient to convert almost all the silver into silver chloride.The points 6-8 refer to the situation where the oxidation time was not sufficient to convert all the silver into silver chloride. Although the zero-current potentials for runs 1-4 are strongly shifted to more positive values, the peaks in the current-potential curves are shifted much less. Nevertheless, the shift can be sufficient to allow reduction of silver chloride by PPD even at a relatively low pH. Fig. 6 shows that at pH 4.5 a mixture current is possible for experiments 1-3, but not for experiment 7. Therefore, at suitable Ag/Au surfaces the formation of silver chloride can be prevented in the presence of PPD as long as the pH is not too low. The influence of the kind and amount of halide ion in solu- tion on the position of the half-wave potentials of the Co1"(NH,),C12+ reduction curves is demonstrated in Fig.7. The full-line curves were recorded at a silver disc and the dashed-line curves at an Ag/Au disc. All sweeps were carried out in the anodic direction, as in Fig. 4. Numbers 2 and 3 refer to experiments in which small concentrations of J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 3 E?12values of 0.5 mmol dm-3 Co"' in the presence of halide at pH 5.6obtained at a silver, a gold and an Ag/Au disc (with 1.00 mC of silver deposited on it) [KCl]/mol dm -[KBr]/mmol dm -[KI]/mmol dm -Ag AU Ag/Au -$202 --40 calomel electrode bromide and iodide were added to the solution. The amount of silver bromide formed depended on the amount of silver deposited on the surface of gold, but it was independent of the concentration of bromide in the solution. The same was true for iodide.Some values of are listed in Table 3. These results can be explained with eqn. (VI) taking uAg< 1 at the Ag/Au disc. The potential was shifted to a more nega- tive value at an Ag/AgBr disc than at an Ag/Au/AgBr disc. Therefore, as with the results obtained with chloride, the reduction of silver bromide occurred more easily on the Ag/Au interface than on the surface of pure silver. As mentioned in the previous section, strong inhibition of the ConI(NH,),Cl2+ reduction occurred at the gold electrode after treating it with a solution of PPD (cf.the current- potential curves G2 and G3 in Fig. 3). At an Ag/Au electrode, however, no inhibition of this kind was observed after pretreating it with N,N-dimethyl-p-phenylenediamine.The curves SG1-SG3 in Fig. 3 are similar to those displayed in Fig. 4 and 5; only the amount of silver chloride formed sub- sequently decreased a little. Inhibition due to the adsorption of PPD was avoided because of the rapid reduction of Co~NH3),C12+ at the silver surface and the simultaneous oxidation of PPD adsorbed at gold. This reaction began immediately after the addition of the cobalt complex and before recording of the current-potential curves. Thus, com- pared with the situation at pure gold discs, a less negative shift of the potential occurred and as a result the mixture current was greater.Some results are summarized in the right-hand column of Table 2. It is clear from the above results that Ag/Au discs display different and catalytically useful properties compared with those of pure silver or gold discs. The introduction of a foreign metal into a metallic catalyst is a well established pro- cedure for modifying surface electronic and geometric struc- tures in order to manipulate catalytic selectivity.' 3*23 Usually the electrocatalytic properties of the bimetallic interfaces are strongly dependent on the relative strengths of the interaction between the reactant and the individual components of the mixed-metal ~urface.~~-~~ Silver and gold have very similar 100 330 110, 150, 350 -10 290 10, 40,250 -50 280 -20,20, 200 -250 150 -240 physical properties (e.g.similar atomic radius and lattice structure).' 5925 The experiments described above have shown that deposition of silver on the surface of gold resulted in a superadditive or synergistic effect. The reduction of Co1yNH3),Cl2+ occurred more rapidly at the surface of silver than at gold while the silver halide, formed according to eqn. (6), could be continuously reduced because PPD could inject electrons into the catalyst at the free surface of gold. The noble metal acting in reactions (1) and (2) is there-fore predominantly silver, whereas the noble metal acting in reaction (7) is gold. The evidence obtained in the present paper thus explains why the inhibition caused by chloride and bromide ions in the catalysis of reactions (1) and (2) by colloidal silver could be overcome by taking a mixture of silver and gold colloids, or better by using Ag/Au particle^.^*^ Conclusion The bimetallic interface consisting of silver islands on a gold substrate possesses unique catalytic properties distinct from those of the pure metals.The formation of silver halide, which is the main autoinhibiting effect in silver-catalysed reactions, can be avoided by electron transfer from the organic reducing agent viu the free gold surface. Thus, even silver halide which has already been formed may be reduced. More- over, the irreversible adsorption of the organic reactant, which is the most important inhibiting effect in gold-catalysed reactions, can be avoided by oxidation of the adsorbed compound via the silver surface.This mutual can- cellation of inhibiting effects has already been used to acceler- ate the colloid-catalysed oxidation of p-phenylenediamines by [Co"'( NH J,Cl] C1 . The authors thank the Deutsche Forschungsgemeinschaft for financial support. References 1 U. Nickel and C-Y. Liu, J. Zmag. Sci., 1990,34,8. 2 R. 0.Farchmin, U. Nickel and M. Spiro, J. Chem. SOC.,Faraday Trans., 1993,89, 229. 3 (a) U. Nickel and C-Y. Liu, J. Photogr. Sci.,1987,35, 191; (b) U. Nickel, C-Y. Liu, P. Lachenmayr and M. Schneider, Bull. SOC. Chim. Fr., 1988,308. 4 M.Spiro, Catal. Today, 1993.17,517. 5 R. 0. Farchmin, Ph.D. Thesis, University of Erlangen-Nurnberg, 1992.6 U. Nickel, K. Kemnitz and W. Jaenicke, J. Chem. SOC., Perkin Trans. 2, 1978, 1188. 7 U.Nickel and W.Jaenicke, J. Chem. SOC.,Perkin Trans. 2, 1980, 1601. 8 Y-H. Chen and U. Nickel, J. Chem. SOC.,Faraday Trans., 1993, 89,2479. 9 U.Nickel, E.Haase and C-Y. Liu, Ber. Bunsenges. Phys. Chem., 1990,94, 726. 10 M.D.Archer and M. Spiro, J. Chem. SOC.A, 1970,82. 11 W.Jaenicke and H. Kobayashi, Elecrrochim. Acta, 1983,28,245. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 623 12 13 14 15 A. M. Creeth and M. Spiro, J. Electroanal. Chem., 1991, 312, 165. (a) C. J. Cali, J. E. Harris, M. E. Bothwell and M. P. Soriaga, Langmuir, 1990 6, 74; (b) M. P. Soriaga, Chem. Rev., 1990, 90, 771. M. Spiro, Chem. SOC. Rev., 1986, 15, 141. (a) J. H. White, M. J. Albarelli and H. D. Abruna, J. Phys. Chem., 1988, 92, 4432; (b) A. Hamelin, J. Chim. Phys., 1991, 88, 1453. 21 22 23 24 M. D. Levi and E. Yu.Pisarevskaya, Electrochim. Acta, 1992,37, 635; G. Horanyi and E. M. Rimayer, J. Electroanal. Chem., 1984,176,339. V. G. Levich, Physicochemical Hydrodynamics, Prentice-Hall, Englewood Cliffs, 1962, p. 69. D. C. Alonzo and B. R. Scharifker, J. Electroanal. Chem., 1989, 274, 167. H. D. Abruna, Electrochemical Interfaces, Modern Techniques for in situ Interface Characterization, VCH Publishers, New 16 17 G. G. Schlessinger, Inorg. Synth., 1967,9, 160. U. Nickel, E. Haase and W. Jaenicke, Ber. Bunsenges. Phys. Chem., 1977,81,849. 25 York, 1991. (a) L-W. H Leung, D. Gosztola and M. J. Weaver, Langmuir, 1987, 3, 45; (b) D. P. Sandoz, R. M. Peekema, H. Freund and 18 19 20 M. Schneider, Ph.D. Thesis, University of Erlangen-Nurnberg, 1989. M. Spiro, J. Chem. Soc., Faruday Trans. I, 1979,75, 1507. M. Noel and K. I. Vasu, Cyclic Voltammetry and the Frontiers of Electrochemistry, Aspect Pub. Ltd., London, 1990. 26 C. F. Morrison Jr., J. Electroanal. Chem., 1970,24, 165. S. Szabo, lnt. Rev. Phys. Chem., 1991,10,207. Paper 3/05888C; Received 30th September, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000617

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(4), 625-627 625 Stability of Thin Polar Films on Non-wettable Substrates Ashutosh Sharma" and Ahmad T. Jameel Department of Chemical Engineering, Indian Institute of Technology at Kanpur, Kanpur 208 016, India The linear and non-linear stabilities of thin (40 nm) apolar and polar films (water, polymers) on non-wettable substrates are correlated to the macroscopic parameters of wetting. The wavelength of instability and the time of film rupture decline as the substrate becomes less wettable by macroscopic drops. The instability of relatively thick (~12nm) completely polar films (eg. water films bounded by octane) evolves very slowly, and the sub-strate may appear to be wettable, notwithstanding a large equilibrium contact angle.In contrast to the apolar films, the non-linear interactions in polar films may remain significant even for the small-amplitude thermal perturbations. Thin (<50 nm) fluid films on non-wettable surfaces are inher- ently unstable'-6 and the growth of interfacial deformations leads to the film breakup and retraction, resulting in drops of finite contact angle. The linear,'-3 as well as the non-linear4*' aspects of the instability have been extensively studied for apolar films (e.g. hydrocarbons) where the excess intermolec- ular energy is derived solely from the apolar Lifshitz-van der Waals (LW) interaction^.^.' However, polar liquids (e.g. water) on the apolar and polar substrates also experience polar interactions variously described as the hydrophobic attraction, hydrogen bonding, 'acid-base' interactions, et~.~.' It is well known' that the interfacial (surface) tensions and the equilibrium wettability (contact angle) of water on any surface can only be described by inclusion of polar 'acid- base' interactions.Indeed, the equilibrium contact angle (wettability) of water, 0, is related to the apolar (nSvLw)and the polar (nSvp)components of the total spreading pressure (ns)by the Young equation cos 0 = 1 + -; ns = nSJ-W + nS,P (1)723 where y23 is the film interfacial tension against the bounding fluid. The term interfacial tension (y,) is reserved for the interface between two condensed media i and j. Ifj is a gas, the term surface tension (yi) is more appropriate for denoting the specific surface energy of material i.nS.' and nSvLware defined in terms of the LW and polar components of various surface (yi) and the interfacial (yij) tensions as' (1 = substrate, 3 = film, 2 = bounding fluid) (24 nS*' = yy2 -y!3 -yp23 = -2y; (if 1 and 2 are apolar) (2b) Thus LW forces promote the non-wettability (film instability) for ns*Lw< 0, which is the case for water films (bounded by air) on low-energy surfaces (e.g. Teflon, aminosilane-treated glass fibre, etc.) with ykW < yiw ( =21.8 mJ m-2 for water).' nSvLwis also negative for water films sandwiched between higher-energy media, i.e. y;", yk" > yf;" also implies nSvLw< 0 from eqn. (2a). Examples are water films on most polymeric and biological substrates bounded by hexadecane.The polar forces also engender non-wettability as nsVpof water on almost any surface is negative due to the large polar cohesive energy of water molecules (75 = 51 mJ m-2).8 The minimum nS*' = -102 mJ m-2 for water is obtained on the apolar substrates (e.g. Teflon), but increased polar inter-actions with the substrate make it less negative [eqn. (2b)l. ns*Lwand nSvpcan be easily evaluated from measurements of contact angles of an apolar liquid and water, in conjunc- tion with eqn. (l).398Briefly, 74" for the substrate can be evaluated with the help of eqn. (1) and (2a) by measuring the contact angle of an apolar hydrocarbon liquid (nS*' = 0) of known surface tension. nSsLwfor water on the same substrate is now obtained from definition (24.Finally, ns*pfor water is determined from eqn. (1) by directly measuring the contact angle on the substrate. In what follows, we investigate the influence of nSvp(and the contact angle) on the stability and kinetics of rupture of water films. Theory The following leading-order non-linear equation describes the evolution of the film thickness, H(x, t)4*5 (3) where t, x and p are time, lateral space coordinate (parallel to the film surface) and the film viscosity, respectively. 4 is the excess energy of the film per unit volume due to intermolecu- lar interactions, which is related to the Gibbs energy per unit area, AG by 4 = aAG/dH, where AG is given by3*798 AG(H) = nSvLW(d;/H2)+ nsyp exp[(do -H)/q (4) where do is the 'cut-off equilibrium distance (=0.158 nm)' where the Born repulsion takes over, and 1 is a correlation length for polar 'acid-base' interactions (I z 0.2-1 nm for water; the best estimate from one study is 0.6 nm).' Clearly, ns,Lwis related to the effective Hamaker constant by the defi- nition, A = -12ndinS*LW,3*8and the change of the Gibbs energy, AG = G(do) -G(m) is as expected given by 2s.Lw + nS,' = ns = y12 -yI3 -y23.Clearly, it is advantageous to write the Gibbs energy of the thin film in the form of eqn. (4), which only involves the macroscopic parameters (nSvLwand nS*') of wetting. As is discussed in the introduction, ns*Lwand nS.' are easily obtained by measurements of equilibrium contact angles of macroscopic drops.In contrast, a direct estimation of the effective Hamaker constant and the strength of polar interactions poses many dificulties, especially for the ill-defined, modified surfaces often encountered in practice. Results and Discussion The linearization of eqn. (3) and (4) around the mean film thickness, h, gives the fastest growing mode of instability of initial amplitude, E, H = h + E sin(2nx//2,)exp(ot), from which a (linear) estimate of the time of rupture' is obtained by setting H = 0, i.e. t, = ln(h/e)/o. The fastest growing (dominant) mode of the linear theory evolves on a length scale (wavelength), A, A: = -(4Z2y23h4/3dz ,s*Lw)( 1 + (h2/doI)2xs9p x exp[(do -h)/fl/6~~.~~} (5)-The linear theory estimate for the minimum time of rupture is t, = by23 h'/3(~~~)~d~](l+ (h2/doO2xp x exp[(do -h)/fl/6xLw}-2 ln(h/&) (6) Clearly, decreased wettability (more negative xLwand x? encourages faster rupture, caused by shorter waves.Results (5) and (6)can also be interpreted in terms of 8 by elimination of xpfrom these equations and eqn. (1). Fig. 1 illustrates the influences of mean film thickness, and the polar spreading pressure (contact angle) on the (linear) time of rupture for a fixed value of xLw= -15 mJ m-2 (water on perfluorolauric acid substrate). For films thinner than 10 nm, polar interaction is significant and a decreased wettability due to the polar interaction (more negative x? causes a greater destabilization of the film.However, for thicker films, a more rapid (exponential) decay of polar forces renders them ineffective and the time of rupture is then con- trolled by the apolar LW forces alone [see eqn. (5) and (6)]. An interesting example in this context is a water film on any substrate bounded by octane, for which ykw = ybw = 21 mJ m-2, and xLw= 0 from eqn. (24. Such completely polar films do not experience excess LW interactions, and relatively thick films should therefore appear to be rather stable, not- withstanding a large (negative) xp and a large equilibrium contact angle, from eqn. (1). The time of rupture (time for appearance of a true three-phase contact zone) for completely polar films is obtained from eqn. (6)by setting xLwto zero, uiz.t, = (%'23/h){l2 exp[(h -d~)/fl/x'h}~In(h/&) (7) For water films bounded by octane, the maximum possible value of I xpI is 102 mJ m-2, from eqn. (2b),and therefore, the minimum time of rupture for water films thicker than 12 nm is more than 30 min. Thus, the contact angle hysteresis in these cases may also be of kinetic origin, which is due to the slow dynamics of the formation of an ultrathin contact zone J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 where the polar interactions take over. In conclusion, for comparable values of the spreading pressure (equilibrium contact angle) relatively thick (>8 nm) films of completely polar systems are considerably more stable than the apolar films. Once a polar film is formed, the surface may remain perfectly wettable for a long time, even though the final equi- librium contact angle, from eqn.(1) is large. However, such thick polar films can still recede and evolve more rapidly into equilibrium drops if local film thickness is decreased (66nm) by external means so that the polar interactions can cause destabilization (receding motion) and a growing microscopic contact zone can form. While the qualitative aspects of the growth of instability are well described by the linear stability [eqn. (5) and (6)],the quantitative predictions require a direct numerical solution of the non-linear equation of evolution (3). The non-linear eqn. (3) was solved by a pseudo-spectral method: Fourier collo- cati~n,~~'~together with periodicity conditions over the wavelength, x E (0,A) and a small amplitude initial dis- turbance, H(x, 0) = h(1 + 0.01 sin 2xx/A).A cosine dis-turbance merely phase-shifts the periodic profile of the film. 24 collocation (grid) points were found satisfactory, and the resulting set of 24 non-linear coupled ordinary differential equations were integrated in time by GEAR algorithm for stiff equations. The length scale (A,,) of the fastest-growing wave, and the corresponding minimum time of rupture t, were determined by solving eqn. (3) for many different values of A in the neighbourhood of AL . The ratios of the non-linear and the linear predictions are shown in Fig. 2. When the polar interactions are important [h < 8 nm, curve (a)], the instability evolves on a length scale (unit cell size), A,,, that may be large compared to the linear estimate, A,.However, for thicker films controlled largely by LW forces, A,, x A,, at least for small initial amplitudes (e.g. thermal perturbations). The size of the unit cell (ca. A,') determines the number of holes per unit area. Experiments with polymer films heated above the glass-transition temperature prove that the equi- librium morphology of the ruptured film is also governed by this factorq6 Further, Fig. 2(b)shows that the non-linearities of both the apolar and polar interactions accelerate the film breakup, i.e., t, < t,. Finally, curve (c) shows the ratio of non-linear times of rupture as obtained for A,, and A,, respectively.For small initial amplitude, t,(A,) x t,(A,), which obviates the need for a 0.2Lit 0 4 a 12 16 0 5 10 15 20 25 h/nm hlnm Fig. 1 Influence of film thickness, nS*' and 8 on the normalized time Fig. 2 Comparison between non-linear and linear theories: (a) of rupture from the linear theory (2vLW = -15 mJ m-'). nssP= (a) 0, A,,/AL; (b)tJtL; (c) tn(An)/tn(AL). (7?s,Lw = -17.17 mJ m-2, nSvp= -60 (b) -5, (c) -10, (6)-30, (e) -60 and (f)-102 mJ m-2. 8 = (a) 37, mJ m-2; symbols represent computed values, lines are only to guide (b)44, (c) 49, (d)68, (e)92 and (f)127 degrees. the eye.) J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 repeated solution of eqn. (3) for different A in order to deter- mine the minimum time of rupture. As is the case for the apolar however, the results may be different for rela- tively large amplitude mechanical perturbations, where non- linear effects are significant from the beginning (work under progress).The theory provides, for the first time, a correlation between the thin-film stability and the macroscopic param- eters of wetting, which should be useful for the design and interpretation of future experiments involving thin polar films (e.g. water, polymers6). References 1 A. Sheludko, Adu. CoIIoid Inteqace Sci., 1967, 1, 391. 2 E. Ruckenstein and R. K. Jain, J. Chem. SOC.,Faraday Trans. 2, 1974,70, 132. 3 A. Sharma, Langmuir, 1993,9, 861, and references therein. 4 M. B. Williams and S. H. Davis, J. Colloid Interfuce Sci., 1982, 90,220. 5 A. Sharma and E. Ruckenstein, J. Colloid Interface Sci., 1986, 113,456. 6 G. Reiter, Phys. Rev. Lett., 1992,68, 75; Langmuir, 1993,9, 1344. 7 J. H. Israelachvili, Intermolecular and Surface Forces, Academic Press, New York, 1985. 8 C. J. van Oss, J. Dispersion Sci. Technol., 1991, 12, 210; C. J. van Oss, M. K. Chaudhury and R. J. Good, Chem. Rev., 1988, 88, 927. 9 C. Canuto, M. Y.Hussaini, A. Quarteroni and T. A. Zang, Spec-tral Methods in Fluid Dynamics, Springer-Verlag, New York, 1988. 10 A. Sharma and A. T. Jameel, Colloid Interface Sci., 1993, 161, 190. Paper 3/06794G ;Received 9th September, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000625

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(4), 629-640 Influence of the Cetyltrimethylammonium Chloride Micellar Pseudophase on the Protolytic Equilibria of Oxyxanthene Dyes at High Bulk Phase Ionic Strength Nikolay 0. Mchedlov-Petrossyan' and Valentina N. Kleshchevnikova Department of Chemistry, Kharkov State University, Kharkov, 310077,Ukraine The acid-base equilibria of fluorescein, sulfonefluorescein, 2,7-dichlorofluorescein, eosin and ethyl eosin have been studied spectrophotometrically in the micellar pseudophase of cetyltrimethylammonium chloride, the charge of micelles being strongly shielded by high counter-ion concentrations in the bulk phase (4.00 mol dm-3 KCI). The 'apparent' ionization constants (K:) and molar absorptivities of all the species have been determined.The completeness of solubilization is confirmed by several approaches, including examination of emission and excitation spectra. Conclusions concerning tautomerism of dye molecules and ions were deduced from absorp- tion spectra; the fractions (a) of tautomers as well as apparent microscopic ionization constants (k")have been evaluated. The medium effects, ApKZ (=pK: -pK:), vary from -1.5 to +2.0and are partly controlled by shifts in the tautomeric equilibria. For the Apk" values (= pk" -pk") the relationship Apk,,,, > ApcC&, is valid. Con- sidering the pKz values at 4.00 mol dm-3 KCI as the first approximation to the 'intrinsic' values (pK!J, and comparing them with the pK,s in aqueous acetone and other solvents, although the influence of micellar micro- environment is qualitatively similar to that of the organic solvents, there are some doubts about the possibility of modelling the values of the whole set of parameters in the pseudophase by that in a 'unique' water-organic compound mixture.Micellar solutions of cationic surfactants are widely used as solvents for various chemical reactions. 'v2 Acid-base proper- ties of substances solubilized by micelles are markedly differ- ent from those in aqueous solution^.^-^ Acid-base indicators are useful probes for the investigation of mi~elles,~*~9~-~since they are generally believed to be located in the Stern layer of ionic miceile~.'*~*~.~.~ The principal parameter measured in such studies is the so-called 'apparent' ionization constant, Ki,,435*8 which is obtained via standard spectrophotometric methods ; the pH values utilized in calculations characterize only the bulk phase and are as a rule measured with a glass electrode.Such systems are actually microheterogeneous ; the dispersed micellized surfactant is considered as a micellar pseudo- phase., The electrostatic approa~h~-~~~*'*-'~ permits the pKP value for the given completely solubilized acid-base indicator couple, HB/B, to be explained in terms of transfer activity coefficients (7) and the potential of Stern layer (Y): YB f"pKf = pK," + log -+ log 2--YF (1)YHB fEB 2.303RT where K," is the thermodynamic ionization constant in water, F is the Faraday constant, R is the gas constant, T is abso- lute temperature and f "are activity coefficients of solubilized species.f" values are not readily available.Since the Stern layer is assumed to be a concentrated salt solution we, as is widely ass~med,~.~.~~*'take fEfEBx 1. The so-called 'intrinsic' ionization constant, Ki , describes the two-phase acid-base equilibria:" PKi = PK: + log(YB/Y,B) (2) The ionization constant in the micellar pseudophase, Kr = (HL)(B,)/(HB), ,is related to Ki by PKF = PK: + log YH (3) Eqn. (3) is often used in attempts to compare the solvent properties of micelles with those of water-organic compound mixture^.^*^* In the case of the micelles of cetyltrimethyl- ammonium chloride (CTAC) and bromide (CTAB), dodecyl- trimethylammonium chloride (DTAC) and bromide (DTAB) as well as cetylpyridinium bromide (CPB) at low bulk phase ionic strength (d0.05mol dm-3) Y > 90 mV.8.'0*13-'6 Th e ApK: values (= pKi -pK,") are usually negative [sometimes even -(3-4) pK, units"].With increased ionic strength of the bulk phase, shielding of the micellar charge occurs, which results in Y decreasing, and the pK: values increasing. The same effect for pK: is explained in the pseudophase ion-exchange mode12.6*10 in terms of the corresponding equilibria shift, e.g.: OH; + Cl; eOH; + Cl, (1) The suffixes m and w denote the micellar and water phases, respectively. The number of counter-ions per micellized sur- factant ion is regarded as constant,2*6*10 but there is evidence that this is not always the case.' At high salt concentrations in the aqueous phase, e.g.4 mol dm-3 Na(K)Cl(Br), the exact residual Y values are, strictly speaking, unknown, but according to some data, they are rather small (e.g.+47 mV for DTAC and + 18 mV for DTAB16). Numerous reports show that under such condi- tions the apparent acidity of solubilized reagents, HB, decreases by ca. 2 orders of magnitude; the pKi values are strongly shifted toward the pKi values.' 'vl' Information con- cerning ionic equilibria in the micellar pseudophase with a (partly) reduced potential is sparse. The present investigation was undertaken to study the pro- tolytic equilibria of fluorescein and its derivatives in the cetyl- trimethylammonium micellar pseudophase at a bulk phase ionic strength of 4.00 mol dm-3 KCl.These dyes, possessing different functional groups (OH, CO,H), ionize in a stepwise manner : H3R+ =H,R + H+ ; K,, (11) H,R=HR- + H'; K,, (111) HR-$RZ-+ H'; K,, (IV) and also undergo tautomeric interconversion (in the case of H,R and, sometimes, HR -). Tautomeric interconversion is quite sensitive to the nature of the solvent and to substituent effects. Previously, we examined in detail the protolytic equi- 630 J. CHEM. SOC. FARADAY TRANS., 1993, VOL. 89 1 100- 100- iI 1 I I I ! ! r c I I 1 I 5 I E II E c I-7 50 E -0 m2--. 15'->1: \ lo\ l/n m 440 460 480 500 520 A/nm Fig. 1 Absorption spectra of fluorescein in CTAC micellar solutions (4 rnol dmP3 KCl); pH = 9.55 (1, spectrum of R2-,7a), 7.94 (2), 7.32 (3), 7.18 (4), 7.02 (9,6.78 (6),6.28 (7), 6.02 (8), 5.86 (9), 5.10 (10); pH, = 0.70 (ll), 0.40 (12), 0.22 (13), -0.40 (14,near to spectrum of H,R+, (la); dotted line: spectrum of HR-(5a), obtained using eqn.(5) libria of fluorescein and its derivatives in in mix- tures of water with acet~ne,~l-~~ and1,4-dio~ane~~~~' DMS0,26*27as well as in micellar solutions of non-ionic surf act ant^^^,^^ and of cetylpyridinium chloride (CPC) at low ionic strength (0.05 mol dmP3 KCl).30-32 Information about the acid-base behaviour of oxyanthene dyes in water in the Fig. 2 Absorption spectra of 2,7-dichlorofluorescein in CTAC micellar solutions (4 mol dm-3 KCI); pH = 9.10 (1, spectrum of R2-, 7b), 7.05 (2), 6.16 (3), 5.95 (4), 5.72 (5), 5.55 (6), 5.25 (7), 5.09 (8)4.89 (9), 4.67 (lo),4.07 (ll), 3.47 (12); dotted line: spectrum of HR-(5b=6b), obtained using eqn.(5) ments were performed at 25.0 & 0.1 "C on a P 363-3 poten- tiometer and pH-121 pH-meter equipped with a ESL-43-07 glass electrode and an Ag/AgCl reference electrode in a cell with liquid junction (3.5 mol dm-3 KCl). Standard buffers (pH 4.01, 6.86 and 9.18) and dilute HCl solutions were used for cell ~alibration.~~ Experimental error did not exceed 0.02 pH units. The dyes were purified as described Cetyltrimethylammonium bromide (Serva, Berlin) was used as received; its purity was determined by two-phase titration withpresence of surfactants is also available in the literat~re.~~-~~ The ionization of lipoidal fluorescein and eosin in a micellar microenvironment l4 as well as the spectra of 2,7-dichloro-fluorescein in reversed anionic micelle~~~ was also described.The aim of the present paper was to obtain the whole set of equilibrium parameters of dyes including the microscopic ionization constants in micellar solutions at high counter-ion concentration and to compare them with the data referring to aqueous and organic media, as well as with those obtained in CPC micellar solutions at low Cl, c~ncentration.~'-~~Such a system (CTAC-4 mol dm-3 KCI) was chosen because the size of C 6H33NMe,f micelles with increasing counter-ion concentration is known to be much smaller in the case of C1-than for Br-.2*37 Experimental Apparatus and Reagents Absorption spectra of dye solutions were measured using SP-46 apparatus (of USSR origin).Emission and excitation spectra were obtained with a Hitachi F-4010. pH measure- a standard solution of sodium lauryl sulfate (H,O-CHCl, , using erythrosin as indicator). Suitable pH values were provided with analytical-grade reagents : hydro-chloric and acetic acid, sodium hydrophosphate and borax ; the standard sodium hydroxide solution was prepared using C0,-free water. Potassium chloride of high quality was further purified by recrystallization. Procedure Measurements were carried out at surfactant concentrations c, = 3 x loP3rnol dm-3. The completeness of CTAB trans- formation into CTAC in 4 mol dmP3 KCl solution results from the value of the ion-exchange constant, Kt;: Cl; + Br; eC1; + Br, ; Kt; (V) According to the various data including those obtained at high ionic strength the K$ values are within the interval 2.7- 5.0.39*40The c.m.c. of CTAC, (1-1.8) x mol dm-3 in water, decreases by more than an order of magnitude at the ionic strength ~hosen.~'.~~ While calculating the quantity of KCl needed for a total ionic strength of 4.00 mol dm-3, the J. CHEM.SOC. FARADAY TRANS., 1994, VOL. 90 63 1 100 r I E c I-0 E m 50 m0 F.. PH PH Fig. 3 Absorption of the dyes in CTAC micellar solutions (4 rnol dm-j KCI) us. pH. (a)fluorescein, 450 nm (1); fluorescein, 500 nm (2); sulfonefluorescein,450 nm (3); sulfonefluorescein, 5 10 nm (4);(b)2,7-dichlorofluorescein,5 15 nm (1); 2,7-dichlorofluorescein,535 nm (2); 2,7-dichlorofluorescein,E (530 nm) + E (535 nm) instead of E (3); eosin, 525 nm (4); eosin, 545 nm (5).ionic strength resulting from the buffer mixtures (acetate, phosphate, borate), not exceeding 0.04 mol dm-3, was taken into account. Hydrochloric acid was used to prepare solu- tions having pH < 3.5, the sum of KCl and HCl concentra- tions being maintained constant (4.00 mol dm-3). Based upon values of ion-exchange constants = 0.50, K;ro4 = 0.78 and K307= 0.7939 the replacement of Cl, by buffer anions is unlikely. In the present study, it was assumed that the pH measurements gave the hydrogen ion activity in the bulk: (H'), = h = In acidic media (pH d 1) the scale pH, = -log cHClwas used in calculations.All experi- ments were carried out at 25 "C. The working solutions were nearly transparent and all spectra were referenced against solvent blanks containing all components except the dyes. The path length was 1-5 cm. During pK: determination concentrations of the dyes were held constant at ca. 8 x mol dmP3 (further lowering leads to the same E values at given pH), but while determin- ing &HrR of fluorescein and 2,7-dichlorofluorescein they were increased to (4-8) x lo-' mol dm-3. There was no indica- tion that the measured results depended upon the increased CTAC concentration. Emission and excitation spectra were measured at dye concentrations of 2 x mol dm-3 (pH 9-1 1).Representative pH-dependences of absorption spectra are depicted in Fig. 1 and 2. The number of different solutions used for pK: determinations was 50, 42, 28, 24 and 13 for fluorescein, 2,7-dichlorofluorescein, eosin, sulfonefluorescein and ethyl eosin, respectively. In the last two cases analytical positions near Amax of intensely coloured species were used for the pK: calculations, while for more complicated equilibria a greater number of wavelengths (15-1 8) were utilized. Typical &-pH dependences are given in Fig. 3. Results Calculation of Ionization Constants and Molar Absorptivities The calculation of pK,, and pK,, of fluorescein from the &-pa: curves in H,O-DMSO mixtures was recently described in The approaches, based on utilizing the whole data at various (with the help of a computer were used for final pK:calculations in the present study.For 2,7- dichlorofluorescein the calculations were carried out analo- gously, except those approaches which are based upon the individual peculiarites of fluorescein The Thamer- Voigt method43 at appropriate A was used to obtain the first approximations of K:, and K:, for further iterations. Differ- ences or sums of E values at various 1 were also used (instead of E) to obtain more marked maxima on the 'titration curves' (Fig. 3). For sulfonefluorescein only equilibria (111) and (IV) were taken into acc~unt,~~,~'-~~ while for ethyl eosin equi- librium (VI) is valid: HReR-+ H+; K,, (VI) Here the calculations were made by the standard procedure. The pKf, value of fluorescein was obtained by treating equi- librium (11) as isolated.The sums of molar absorptivities at 440,445 and 450 nm for solutions from pH, = 0.00 to pH, = 0.70 were used in the standard formula (4) pK:O = pHc + log[(& -&HzR)/(&H~R+ -(4) while the sum of E of the solution with cHCl= 2.5 mol dmP3 and cKC, = 1.5 mol dmW3 at 435, 440 and 445 nm, multiplied by 1.086, was used as &H3R+. Thus corrections were made for the solvatochromic shift of the H3R+ band (see below) and incomplete transformation of the dye into the cationic form at pH, = -0.4 (E,,, = 50 x lo3, while in water at pH, d 0 E,,, = 54.3 x lo3 dm3 mol- cm-').The results are presented in Table 1 together with the pKr values. The pK: values in CPC micelles (0.05 mol dm-3 KCl), obtained previo~sly,~~-~~ are also tabulated. Direct compari- son of our results with those obtained for octadecanoyl aminofluorescein (OAF) and hexadecanoyl aminoeosin (HAE)14 is impeded. In DTAC micelles (c, = 0.05 mol dm-3, without salt additives): pKt, = 5.78, pKf, = 3.71 (+0.05), pK;, < 0.5 (OAF) and pKE, = 2.87 k0.05, pK:, = 0.8 L-0.2 (HAE).l4 In DTAB micelles (4 mol dm-3 NaBr) the pK: values are reported as 7.47, 5.66 and 0.82 (k0.07) for OAF and 4.53, 2.32 (k0.05) for HAE.14 The pKr values in bromide systems at ionic strength in the bulk 4 mol dm-3 solution are 632 J. CHEM.SOC. FARADAY TRANS., 1994, VOL. 90 Table 1 The apparent ionization constants of oxyxanthenes in micellar (3 x rnol dm-3) solutions of cationic surfactants and the ther- modynamic pK, values in aqueous solutions ~~~ ~ CPC micelles, CTAC micelles, water 0.05 rnol dmP3 KC1" 4.00 rnol dm-3 KCl dye PKZO Pq PKZ2 PK:o PK:, PK:, PK:o PK:, PK:, fluorescein 2.14' 4.45' 6.80' 0.98 3.60 5.54 0.60 k0.10' 6.41 & 0.10 7.17 f0.06 sulfonefluorescein -3.22d 6.76d -0.9 5.46 -2.33 k0.05 7.00 f0.012,7-dichlorofluorescein 0.35' 4.00" 5.19' -3.58 3.70 < -03 5.50 f0.05 5.79 f0.06 eosin -2.@ 2.818 3.7Y -0.5' 2.82 -1.83 f0.07h 5.76 f0.06'' ethyl eosin -1.91' --0.V --1.11 & 0.03' -" From ref. 30-32; 20 "C.'From ref. 19; 25 "C. 'In the scale pH, = -log cHC,.From ref. 30-32; 20 "C. From ref. 21 ; 25 "C. f In the H, scale; -2.04 f0.03 (H,O-H,SO,), -1.98 k0.12 (H,O-HCI), -1.83 f0.09 (H,O-HCIO,); from ref. 42. From ref. 20; 20°C. In the pre- liminary comm~nication~~ the values 1.81 k0.11 and 5.89 k0.11 were reported as obtained from a smaller number of experimental data. Obtained by extrapolating the dependence of pK,, on 1/D in H,O-Me,CO mixtures; from ref. 23. j From ref. 31. known to be CQ. 0.5 units higher than those in chloride should be carried out: systems.l3.1 5.16.18.44 Spectral characteristics of R2-ions are measured directly at pH 9.5-11.5 and for R- of ethyl eosin at pH 6-7 (because of the C02Et group hydrolysis in aqueous alkali). The molar The H,R spectra are given in Fig. 5. Comparing them against absorptivities of HR- ions were obtained together with Kf, experimental curves (dotted lines) shows the necessity of the and Kf2 and then recalculated within the whole visible region use of eqn. (6).by using the pH area of maximal yield of the sought species Spectral characteristics of all the species differ from those in water (Tables 2 and 3), but are similar to those obtained in in solutions (pKf, < pH < pKf,): the presence of CPC micelles (0.05 mol dm-3 KCl).30332 The cases with the spectra of the H3R+ ion of fluorescein and H2R of sulfonefluorescein are more complicated. While for eosin there exists a narrow pH interval near pH ca. 4.0 where the directly measured E values are close to eHR-(Fig. 4), and for sulfonefluorescein this interval is wider, in the case of fluorescein and its dichloro derivative the HR--spectra, obtained by eqn.(5) (Fig. 1 and 2: dotted lines), do 4.5 not coincide with any spectra, measured directly. The stan- dard deviations of EHR-values are &(5-7)%. Because of weak intensities of H2R species of fluorescein and 2,7-dichlorofluorescein exact numerical &HzR values are 4.0 -not needed for pKf calculations. Traces of intensely coloured species, however, may strongly influence the spectra mea- sured in the pH region of H2R predominance. In order to 'remove' the ion contribution the following &HzR verification 3.5. , 2 .-I E c I-2.5.E m E U m z---. 40-450 500 550 ..-Alnm I ..,, I,--, 500 550 Fig.5 Absorption spectra of neutral forms of the dyes in CTAC A/n m micellar solutions (4 rnol dm-3 KCI). (1)-(3): H,R (3e4); (1) fluo- rescein, (2) 2,7-dichlorofluorescein, (3) eosin; (4) H,R of sul-Fig. 4 Absorption spectra of anions in CTAC micellar solutions (4 fonefluorescein (8); (1') and (2') spectra of fluorescein and 2,7-rnol dm-3 KCI): (1) R2- of sulfonefluorescein (lo), (2) RZ-of eosin dichlorofluorescein at pH 3.47; (3') spectrum of eosin (cHC,= 1.1 rnol (7c), (3) HR-of eosin (k),obtained using eqn. (5), (4) R-of ethyl dm-3); (5) HR of ethyl eosin in 2.5 rnol dm-3 HCI (11).(6) HR- of eosin (12) sulfonefluorescein(9). J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 2 Spectral characteristics of various species in micellar solutions and in water AmJnm(~JlO3 dm3 mol-' cm-') CTAC micelles, dye, species 4.00 mol dm -KCl H20a fluorescein H3R+ (la) 444-445 437 (54.3) H2R (2a e3a F? 4a) 465 (0.393), 490 (0.361) 437 (13.9), 470-485 (4-3.1) HR-(5a) 455 (24.9), 480 (24.0) 454-474 (32.7-33.8) R2-(7a) 500 (85.9) 491 (88.0) sulfonefluorescein H2R (8) 452 (56.6) 440 (52.2)HR-(9) 460 (29.9), 490 (25.5) 455, 480 (29) R2-(10) 512 (83.7) 495 (84.0) 2,7-dichlorofluorescein H2R (3b +4b) 470 (0.822), 500 (0.766) 460 (8.96), 485 (8.70) HR -(5be6b) 525 (57.4) 465-470 (24.1), 490 (28.3) R2-(7b) 515 (80.1) 502 (75.02) eosin H,R (3c F? 4c) 480 (7.92) 480-485 (8.5)* HR-(6c) 538-540 (97.6) 517-519 (81.9) R2-(7c) 526.5 (103.4) 5 15 (96.7) ethyl eosin CHR (11) 485.5 (21.6) -R-(12) 542.5 (97.1) 520 (97.3)' a As reported in ref. 19-21,42.'pH 0-0.5 (H2S0,). Low solubility. 5% EtOH or Me2C0. Location of Dye Species in the Microheterogeneous System The HR -spectrum of 2,7-dichlorofluorescein (Fig. 2, Table The bathochromic shifts (9-17 nm) of R2-absorption bands 2) differs decisively from that in water (where its configu- as compared to the spectra in water (Table 2) are usual for ration is similar to that of the HR-spectrum of these dyes in the case of replacing H20 by an organic fluorescein2'). Similar changes in the spectra of monoanionic The same is true for the HR- ion of eosin 2,7-dichlorofluorescein in water-acetone mixtures have been in CTAC micelles the shifts shown to be connected with the shift of equilibria from car- and R- ion of ethyl eosin;23,25*26 are 21-22 nm. For the HR- ions of fluorescein and sul- boxylate tautomer towards the phenolate one.21 fonefluorescein the solvatochromic shifts are less dramatic, The bathochromic displacement of the H2R band of sul- but they are still registered upon transferring the species from fonefluorescein (Table 2) is appreciable (12 nm); for this dye water to micelles of cationic surfactants.The closeness of A,,, the formation of colourless tautomers even in pure organic and in micellar solutions the values of R2- bands of fluorescein and eosin and HR- of solvents is not specific,21*26,27*42 eosin in micellar media, obtained in the present study (Table E,,, value is even higher than in water (Fig.5, Table 2). 2), to those of the lipoidal (solubilized) analogues under All the mentioned spectral effects, occurring in micellar similar conditions'" also confirms the completeness of solu- media (as compared to aqueous solutions), as well as pK: bilization. The deviation of the A,,(HR-) value of OAF values, are almost constant with a tenfold or more decrease (494-496 nm'") from our value for fluorescein (Table 2) may in the ratio cdye/c,. These observations are evidence of soh-be caused by the difference in assigning bands because the bilization of the species studied. If the concentrations of solu- HR- spectrum was not measured directly, see eqn. (5). bilized and non-solubilized species are commensurable, then The decrease in intensity of H2R spectra for fluorescein a shift of the distribution equilibria should be observed and 2,7-dichlorofluorescein (Fig.5), as compared to those in during the variations in the number of micellized surfactant aqueous media (Table 2), is attributed to shifts in the tauto- molecules (ions) per dye molecule. meric equilibria. The solubility of neutral forms of eosin In Table 3 the A,, values of fluorescence and excitation (H2R) and ethyl eosin (HR) increases essentially in the pres- spectra of the most hydrophilic species (R2-) are presented. ence of micelles as compared to water, where solubility is The data prove that the dianions R2- are completely solu- extremely low (ca. 2 x mol dm-3 and <loF6 mol bilized. The A,, of excitation are independent of the analyti- dm-3, re~pectively~'*"~). cal wavelengths used, i.e.no traces of unsolubilized (hydrated Table 3 The characteristic fluorescence and absorption (A,.Jnm) of R2- ions in water, aqueous KCl and in micellar (3 x mol dm-3) solutions of cationic surfactants emission excitation absorption CTAC, CTAC, CTAC, 4 rnol d m-3 4 mol dm-3 4 rnol d mP3 4 mol dmP3 4 rnol dm-3 dye H20 KCl CPC KCl H,O KCl CPC KCl H20 CPC KCl fluorescein 515 518 530 526 492 492 505 500 491 504.5 500 sulfonefluorescein 518 52 1 535 533 496 500 514 512 495 512 512 2,7-dichlorofluorescein 525 526 538 537 504 504 517 516 502 514.5 515 eosin 537 540 548 548 514 515 528 527 515 527 526.5 in the bulk phase) dye ions are revealed.The legitimacy of emission and excitation spectra used for such purposes (without trying to obtain information on peculiarities of solu- bilization of bound dyes in the pseudophase) is based on the much shorter fluorescence lifetime of the xanthene dyes (R2-)45-47 than the lifetime of micelle~.~' At pH, 2 0 I,,,(H3R+) of fluorescein ($44-445 nm) differs markedly from that in water [437 nm in HCl solutions (pH 0 ~~-~~to l), as well as at c~ = 4.0 mol dm-3], which allows one to consider the cationic species to be solubilized. For the OAF cation in DTAB micelles (4 mol dm-3 NaBr) I,, = 448nm.I4 At pH, < 0 the superseding of H3R+ of fluorescein from CTAC micellar surface or some other effects of HCl are possible [A,,,(H3R+) = 439-440 nm].In the system CPC-0.05 mol dm-3 KC1-HC1, 4,,,,(H3R+) is at 437 nm, as in water. Apparently the cation H3R+ is unsolubilized at high Y values; under this condition the pKf, lowering (=0.98; pKro = 2.14) occurs only because of H2R solubilization for the completely solubilized OAF in DTAC micelles at low [Cl-1,: pKi,(0.514. The other species of fluo-rescein and its derivatives, judging by the I,,, values and molar absorptivities as well as solubilities of neutral forms, are solubilized in CPC rni~elles.~'-~~.~~ The bathochromic shift of the H2R band of sulfonefluorescein in CPC micelles (5 nm) is less than in the micellar pseudophase with reduced Y value (12 nm), but this may be connected with the pecu- J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 liarities in the position of cationic chromophores within the micelles of both types. Discussion Detailed Scheme of Protolytic Equilibria in the Micellar Pseudo phase ~ Scheme 1, describing the interconversions of protolytic species (1 -, . . -,7), was discussed in previous pub-lication~.'~-~~'~~-~~The main assumptions, used in the cal- culation of the fractions (a) of various tautomers of H2R and HR-, are: (i) The lactone 4 is colourless and (ii) the ioniza- tion of the carboxylic group exercises a relatively slight influ- ence on the absorption band of the xanthene chromophore in the visible region. Hence the spectra of species 2, 3 and 6 for the dye with the given substituents in the positions 2, 4, 5, 7 or/and in the phthalic residue, are close to those of 1,5 and 7 respectively; in the last case the chromophore is most sensi- tive to 2'-substituents, and the band of 6 is red-shifted (ca.3-15 nm) compared to that for 7.20-26 The spectra (Fig. 1 and 4) show that the HR- ion of fluo-rescein exists as Sa, while the HR- ion of eosin exists as 6c. This is proved by comparing the spectra of these two mono- anionic dyes with the spectra of the HR- form of sul-fonefluorescein (9, Scheme 2) and of R- of ethyl eosin (12, IbC02-I -b""; 7 Scheme 1 Protolytic equilibria of oxyxanthenes. la-7a: unsubstituted fluorescein, lb7b: 2,7-dichlorofluorescein, lc-7c: 2,4,5,7-tetra-bromofluorescein (eosin) J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 635 8 9 10 Scheme 2 Ionization of sulfonefluorescein Br Br t3r Br C02Et C02Etty 0 11 12 Scheme 3 Ionization of ethyl eosin Scheme 3) respectively (Fig. 4 and 5). The replacement of the CO, group by SO, results in an additional red shift of the R2-ion band2' (Fig. 1 and 4, Tables 2 and 3). The HR-absorption intensity of eosin is somewhat lower than that of R2- [&,,(HR-) = 0.94 cmaX(R2-)], but this ratio remains almost constant (ca. 0.9) in various organic sol-yentS,20-26.29-32.42 which permits it to be regarded as a manifestation of 6c chromophore properties rather than a sign of the presence of 5c (see below). In the case of the HR-ion of 2,7-dichlorofluorescein in water the 5b tautomer pre- dominates,2' which may explain its incomplete transform- ation into 6b in micellar media.For calculating asb and a6b 2 = 525 nm was used: Taking &,,(6b) equal to &,,(7b) = &,,(R2-) [actually reali- zing the correction for the shift in bands, caused by 2'-substituent replacement (CO, 4 C02H)] and neglecting the 5b absorption in this region one can obtain ash = 0.72. Pro- ceding from the ~,,,,,(6)/~~~~(7) of 2,4,5,7-tetra-ratio halogenoderivatives in cationic micelles (average value : 0.9030-32.42 ) and assuming that the &Sb/&6bratio at Amax(6b)is equal to ratio of ethyl eosin at 542.5 nm (=0.082, Fig. 4 and 5), we have calculated the values a5b= 0.222 and a6b = 0.778, which we consider most reliable. Then KTx= [6b]/[5b] = 3.5.Similarly in CPC micelles (0.05 mol dm-3 KCl), in these for &,,(R2 -) of 2,7-dichloro- fluorescein in H20 the mean value of 101 x lo3 (503 nm)48 and 97 x lo3 dm3 mol-' cm-' (502.5 nm)49 was postulated and the apparent &,,(R2-) value in micellar solutions, reported previo~sly,~~,~~ is also higher than in Table 2. For the spectra of neutral forms (Fig. 5) the following equation is derived : 'H2R = '2 '2 + '3 '3 (9) The H2R spectra of fluorescein does not contain obvious fea- tures of 2a (which must absorb as la and 8, Fig. 1 and 5 and Table 2); the coexistence of 3a [absorbing as 5a (Fig. 1) and 9 (Fig. 5)] with the colourless lactone 4a is evident. In the case of eosin, the zwitterionic tautomer 2c is highly improbable because of the strong acidity of the oxy groups.The coloured fraction of H2R (3c) has a spectrum analogous to that of HR (11) of ethyl eosin (Fig. 5). The a3 values were calculated as : '3 = &H2dE3 (10) from the data near AmaX(H2R).The value e3, = 27.4 x lo3 dm3 mol-' cm-' [average of E,,, of 5a and 91 was used to offind afa = 0.014, a4' = 0.986. ~,,,~,(ll) ethyl eosin, 21.6 x lo3 dm3 mol-' cm-', was used in the case of eosin: a3, = 0.37, a4, = 0.63. For the dihalogeno-derivative the uti- lization of the average value 24.5 x lo3 dm3 mol-' cm-' leads to a3b = 0.034, a4b = 0.966. Thus the KT (tautomerisation constant) values, KT = [4]/[3], are equal to 70.4, 28.4 and 1.70 for fluorescein, 2,7-dichlorofluorescein and eosin, respectively. Apparent Microscopic Ionization Constants, k' It is easy to show that: pKfO = pkt,OH + log a3 = pk", C02H + log a2 (l pKfl = pk?,COzH -log '3 + log 'S = pk",oH -log a3 + log 66 = pkt, -log a2 + log a5 (12) pKf2 = pk?,OH -log '5 = pk", CO2H -log a6 (13) where k" are microscopic ionization constants (actually 'apparent' microscopic ionization constants): k", C02H = hc21/[11, k", OH = hC31/C11, kt,Z = h[51/c21, k",COzH = hc51/~31, k", OH = h[61/[31, k?, OH = hcwc5i, k;.CO2H = h[7]/[6], see Scheme 1. The pk" values are presented in Table 4. The fraction of the zwitterion aZn was undetectable by means of the absorption spectra (see above). However, it can be shown that: 10g(a3a/a2a)= pk", C02H -pk\, OH = pk?,C02H -pkt,Z (14) Using the pk;,, value of sulfonefluorescein as the corre-sponding value for fluorescein and assuming that pk", C02H is not more than 1 unit lower than pk",,,,, (see below), it is easy to evaluate: a3JaZp x lo2-lo3, which agrees with the spectral data.The pk;, C02H value of eosin (= pKi2) is higher than pk", COzH of fluorescein. The reason may be the influence of the negative charge delocalized in the xanthene moiety of 6c. Such an influence (and, in turn, the influence of the charge J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 4 Apparent microscopic ionization constants (P)of the oxyxanthene series in CTAC micellar pseudophase (4.00 mol dmF3 KC1 in the bulk) and the medium effects (ApP, AApP) fluorescein, k,,, OH sulfonefluorescein, k eosin, k 1, OH ethyl eosin, k 1, OH 2,7-dichlorofluorescein,k OH fluorescein, k2, OH sulfonefluorescein, k2,OH 2,7-dichlorofluorescein, k2,OH 2,7-dichlorofluorescein,k 1, mzH fluorescein, k 1.C02H 2,7-dichlorofluorescein,k2, COzH eosin, k2, COzH 2.45 3.10 -0.65 0.1" 2.33 3.22 -0.9 1.4 1.40 2.40 -1.0 0.9' 1.11 4.14 (1.9) - ( -0.8) - 1.1' 1.3 7.17 6.80 0.37 1.6 7.00 6.76 0.24 1.5 (5.14) 5.19 (0) (1.8) (4.68) 4.56 3.5 3.49 (1.2)1.1 (1.7)2.4 5.68 - - 2.2 5.76 3.75 2.0 2.9 ~~~~ The cationic species are probably desolubilized in CPC micelles (0.05 mol dm-3 KCl). The pk.,, OH of eosin and ethyl eosin in CPC micelles are obtained in the pH, scale. in the phthalic residue on the acidity of OH groups) may be qualitatively expressed through the Bjerrum-Kirkwood-Westheimer equation :50 6 = 241/Deffp where 6 is the additional rise in pk of the ionizing group, p/A is the distance between the latter and the negatively charged group, Deff is the 'effective' relative permittivity.If the xanth- ene moiety is positively charged (la species) the pk of the carboxylic group (pk,, C02H) can be expected to be lower than pk~,,,,,. In the media under study the pk:,, of sulfone- fluorescein is even less than pkt, OH of fluorescein, against eqn. (15) and contrary to H,O-DMSO, where pk,,, is 0.12-0.90 units higher than pk,,oH.27 The reasons may be the follow- ing: (i) pKf, ,used for pk", OH determination, is calculated on the pH, scale (perhaps at high KCl concentrations f; > 1); (ii) the peculiarities of the arrangement of solubilized cations la may be significant; in CTAC micelles these species may be shifted toward the water phase.The pk;,,, of eosin satisfactorily coincides with pKr, (=pk", OH) of ethyl eosin (Scheme 3), the latter value being 0.2-0.5 lower than the former in various organic solvents [e.g., see Fig. qb)]. While the pkt, C02H and pki, COzH values of 2,7-dichloro- fluorescein are close to the corresponding values of fluores- cein and eosin, respectively, the pk:,,, and pki,,, of the dihalogeno derivative are essentially higher than pk:, OH of eosin and ethyl eosin, but essentially lower than pk;,,, of fluoresecein and sulfonefluorescein (Table 4).This is undoubtedly the consequence of substitution; the pki, OH values of phenolsulfonephthalein and its tetrabromo-derivative in CTAC micellar pseudophase (4.00 mol dm- KC1 in the bulk) are 8.71 and 3.86, respectively.'* From Scheme 1 the following equation can be derived: Using the pk;,,,, pki,C02H values for eosin and pk;,,,,,, pk;,,, values for fluorescein and taking into account eqn. (15) KTxcan be calculated: KTx > lo3 (eosin) and KTx x 0.01 (fluorescein). The results agree with the absorption spectra of the HR -forms. Taking into account the degree of precision in confirming the assumptions used in spectra modelling of single tauto- mers as well as the errors of E values, used in a calculations, the a errors can reach ca.20%. Hence the pk" standard devi- ations are on the whole greater than those of pKr. The least exact is the aSb value (to emphasize this fact the values pk;, C02H and pki, OH of 2,7-dichlorofluorescein are given in Table 4 in parenthesis). If + a6b, then pk?,,, = 4.03, pk", C02H = 5.79. The Stern layer is in reality a region of high electrolyte concentration (ca. 3-7 mol dm-3).6,10 It should be noted that the salt effects on the pK,,, pK,, and pK,, of fluorescein in the H,O-KCl system (to 2.3 mol dm-3 KC1) are of standard nature (the values vary within the ranges 2.14-2.40,4.45-4.18 and 6.80-6.26 respectively), and aza, as, and a,. values are not influenced by cKcI.42That is why the effects observed in CTAC micelles cannot by caused by the electrolytic environ- ment itself. Medium Effects for Macroscopic and Microscopic Apparent Ionization Constants (ApK ,AN') The ApKr [k(0.07-0.191 in CTAC micellar solutions vary within the range of 3.4 units: from -1.4 (fluorescein, ApKt,), -0.98 (eosin, ApKf,) and -0.89 (sulfonefluorescein, ApKf,) to 2.01 (eosin, ApK;,) and 1.96 (fluorescein, ApKf,).The ApKr values can be devided into Apk" (=pk" -pk") and A log a [eqn. (11-13)]. The Apk" values are given in Table 4; though they are relatively small, the differences in medium effects are subjected to the same regularities as in organic solvents2'-' and aqueous solutions of non-ionic polymer^,^' e.g. for fluorescein and sulfonefluorescein : OH (and 'Pk",, Z) < 'Pk;, OH < 'Pk?, COzH (17) Here, as usua1,38*50*52 the medium effects for carboxylic acids are higher than for phenols.The charge types of cationic and zwitterionic acids are known to cause even the decrease in their pK, values (here pkt,,, and pk",,,) on addition of organic co-solven ts.' 9 38*5' The value for eosin is somewhat lower again (Table 4), but this is perhaps the result of the differentiating impact of the micellar environment. Apki, C02H (eosin) is higher than Apk", C02H (fluorescein), as is usual for anionic and neutral acids with the same ionizing gro~p.~' Medium Effects in Aqueous Acetone Here the ApKf and Apk" values are compared with the medium effects in the H,O-Me,CO system; it does not mean that we regard the solvent properties of these mixtures as being the closest to those of micellar pseudophase.The pK, values compiled from previous studies' 1-23 are plotted against acetone concentrations (wt.%) (Fig. 6). The thermody- J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Jf~'""'''20 40 60 80 J\ -' 20 ' 40 ' 60 ' 80 ' 20 40 60 80 wt.% acetone Fig. 6 Dependences of pK, on wt.% of acetone. (a)fluorescein: (1) pK,, ,(2) pK,, ,(3) pK,, ( =pk2, OH)r (4) pk,, OH, (5) pk,, COIH. (6)eosin: (1) pK,, ,(2) pK,, (= pk,, C02H), (3) pk,. OH;(4)pK,, (= pk,, OH) of ethyl eosin. (c) 2,7-dichlorofluorescein:(1) pK,, ,(2) pK,, . namic pK, values (usually kO.05) relate to pa; scales. For and pK,, values of fluorescein and 2,7-dichlorofluorescein fluorescein the same trends as in H,O-DMSO mixtures27 are became closer, while pK,, and pK,, values of eosin and pK,, evident.HR- exist as 5a, the 2a tautomer disappears: and pK,, values of fluoresecein grew further apart. In all the cases a3a< a3b < age.APKao = AP~,,OH + A log ~3n (18) ApKal = APkl,C02H-A log u3a (19) The Properties of the CTAC Micellar Pseudophase at High ApKa2 = Apk2, OH (20) Ionic Strength of the Bulk Phase as compared with Water-rganic Compound Mixtures The sharp decrease in aSa(from 0.111 in H20 to 5 x in 90% Me,CO), U-shaped pk,, OH dependence on Me2C0 The pKr values [see eqn. (3)] correspond to the pK, in content, typical for cationic acid^,^^.^^ and more marked H20-Me2C0 at certain unknown acetone content. Even if medium effects for carboxylic groups than for oxy pKt at 4 mol dmP3 KCl in the bulk phase can be considered groups 2 77387 50.5 2 as an approximation of pKa (neglecting the 'residual' quan- tity YF/2.303RT), the log yH value [see eqn.(3)] is unknown ApkO,OH < Apk2,OH < APkl,CO2H (21) (essentially ex trathermod ynamic), and no attempts were 7, made here to evaluate it. However, if the and Y values are lead to different changes of Kao/K,, and K,,/K,, ratios: the same for all the solubilized dyes, it could be expected that the whole set of parameters (at least pk" values) would corre- spond to one fixed Me,CO content. COzH -AP~,A(PKa1 -PKaJ = AP~,, OH -2A log ~3a (22) This is, however, not the case. From the curves in Fig. 6 (plotted on a large scale) it is found that the 'apparent' pK:,, A(pKa2 -PKa1) = A~k2,OH -AP~,,C02H + A log a3a pKf2 (= pk", OH) a;nd pkt, C02H of fluorescein in CTAC micel- (23) les (4 mol dm- KCl) correspond to 27, 19.4 and 21% Me2C0, respectively, pKf , and pK:, of 2,7-dichloro-As a result K,,/K,, increases from 204 in H,O to 3.8 x 10" fluorescein and pK:2 ( =pk", C02H) of eosin correspond to 16, 22 and 28.5% Me,CO, respectively.In the cases of pK:, andin 90% Me,CO, while K,,/Ka2 decreases from 224 in H,O to 0.05 (inversion of ionization constants) in 90% Me2C0. In contrast, for eosin HR- exists as 6c and aJCdoes not decrease as sharply (a3c= 0.057 in 90% Me2CO): pk:,,, of fluorescein, pK:, and pk;.,, of eosin and pK:, (=pk;,,,) of ethyl eosin all the values are lower than corre- sponding pK, values in H20-Me,CO mixtures.In H20-EtOH42 the K,, , K,, and kl,C02H of fluorescein, equal to the apparent (micellar) values, are measured at, respectively, 44-45, 25 and 25 wt.% of EtOH, pK,, = 5.76 of A(pKa2 -PKa,) = Apk2, cO~H-AP~,,OH + A log a3c (24) As pk2,co2H rises more sharply than pkl,OH, K,,/Ka2 increases from 8.7 in H,O to 4 x lo3in 90% Me,CO. In the case of 2,7-dichlorofluorescein the interpretation is more complicated because, the shift not only from 3b to 4b (a3b = 0.010 in 90% Me,CO), but also from 5b to 6b occurs. As a result K,,/K,, changes from 15.5 in H,O to 1.3 in 90% Me,CO. All the trends observed for the dyes under discussion in aqueous acetone as well as in other solvents (Table 5) qualit- atively followed those stated in the present paper for the CTAC micellar pseudophase (4 mol dmP3 KCl); the pK,, eosin corresponds to 45% EtOH.However, the values, equal to pKto and pkt, OH of fluorescein, pKr, and pk", OH of eosin and pK:, (= pk;, OH) of ethyl eosin, are below the correspond- ing pK, values within the whole range of EtOH content.42 While the k~,,,,, and k:,,, values of fluorescein corre- spond to the k values in ca. 22-42% DMS0,27 the values, equal to K:,, K:, and Kr2 (=k;,,,), were recorded in CQ. 50% DMS0.27 In the H20-1,4-dioxane system4, the values K,, and K,, of fluorescein, equal to K:, and K:,, are obtained at 27-30% of the organic co-solvent; k1,C02H, equal J. CHEM. SOC.FARADAY TRANS., 1994, VOL. 90 Table 5 The values of log(Ka,/Ka2) and of the quinoid tautomer fractions (a3)for fluorescein, 2,7-dichlorofluorescein and eosin in various media lodKailKa2) = ~Ka2-PKai a3 2,7-dichloro- 2,7-dichloro- media fluorescein fluorescein eosin fluorescein fluorescein eosin H,O" 21 wt.% ficolld 2.35 1.04 1.19 - 0.94 1.64 O.lllb 0.0823' 0.32 - 0.39' 0.21 16 wt.% Me2COf 1.40 0.5 2.06 0.072 0.10 0.40 CPC micellese 1.94 0.12 2.32 0.038 0.108 -h CTAC micelles' 0.76 0.29 3.93 0.014 0.034 0.37 64wt.% l,4dioxaneJ 0.1 0.19 3.33 5.7 x 10-3 0.022 0.09 91 wt.% DMSO~ -1.35 0.07 3.91 1.7 x 10-3 0.014 0.1 From ref. 19-21 and 42. * aZn= 0.218. 'The cJe value is taken to be 21.6 x lo3 dm3 mol-' cm-'.From ref. 51; ionic strength: 0.05 rnol dm-3 NaCI. aZ. = 0.0287. f From ref. 21-23, 42. From ref. 30, 32; ionic strength: 0.05 rnol dm-3 KCI. The value cmaX(H,R) = 21 x lo3 dm3 mol-' cm-' is calculated from equilibrium data in the acidic pH, region. '4.00 mol dm-3 KCl. j From ref. 25,42. From ref. 26,27, 42. to kt,,2H, corresponds to 20% 1,4-dioxane, while all pK,, and pk,,oH within the range 0-64% 1,4-dioxane are higher than the values of corresponding apparent (micellar) con- stants. The KT values in micellar medium can be directly com- pared with those in water-rganic compound mixtures without having Y and yH values at our disposal. While KT values, equal to 70 and 28, are determined for fluorescein and 2,7-dichlorofluorescein, respectively, in 36.5% Me,CO, the value K, = 1.7 of eosin corresponds to ca.18% Me2C0. The KT value for the latter dye is less sensitive to variations in en~ironment~~.~~.~~(Table 5). The 'micellar' value for fluo- rescein is given for 40% 1,4-dioxane and 52% DMSO. The D values of these three mixtures lie within the wide range 44.4- 75.5. Much closer are the E: (Dimroth-Reichardt parameter) values (0.78, 0.75 and 0.75 for 36.5% Me,CO, 40% 1,4-dioxane and 52% DMS053), which emphasises the role of the H-bonding ability of the solvents in the determination of the KT value. The E: value for CTAC micelles is not available in the literature; for DTAC micelles Ef =0.74 and 0.70 in aqueous solutions and at cNaCl= 4 mol dm-3.15 The signifi- cance of H-bonds in 3a stabilization stresses the fact that even in pure EtOH (D = 25, EF =0.66) the K, value (= 58)42 is lower than in CTAC micelles or in mixtures of water with 'aprotic' solvents with ET =0.75-0.78. The differences of Apk" values for stepwise ionization of fluorescein [eqn.(25)] and of eosin [eqn. (26)] may be expressed as 'Pk;, OH -'Pk;, CO2H = log Y7n +log Y3n -2 log ysn = -0.7 (25) The decrease in values on introducing four bromine atoms occurs for all the eosin species and to some extent this effect is mutually compensated in eqn. (26). Therefore, the essential difference in the above log y combinations for fluorescein and eosin reflects more strongly the stabilization of the anion 6c with delocalized charge as compared to the carboxylate anion (5a) in micellar pseudophase. This effect, being of the same nature as those described by expressions (17) and (21), is typical for the transfer from water to media with lower H- bonding ability.' 7s2s4 The modern structural models of micelles of the type discussed' ,2*46*5 presume the entry of hydrocarbon chains into the Stern layer (or rather 'Stern region'').Such proxim- ity of -CH2- groups to the dyes examined hinders H- bonding with water molecules, causing the effects described above. Such a mixture of water, electrolyte (3-7 mol dm-3) and hydrocarbon is unattainable in homogeneous systems, and this is the reason for difficulties in modelling the Stern region effect on the protolytic equilibria through comparing with medium effects in aqueous acetone, 1,4-dioxane, ethanol, DMSO etc.The location of different dyespecies within the Stern region, in particular the depth of penetration, can also contribute to the resulting medium effects (ApKf:, Apk", A log a), but the simple model used in the present study presup- poses an identical microenvironment [eqn. (1)-(311 with a fixed set of parameters of the pseudophase for all the particles solubilized. Transfer from CPC Micelles, [C1-] =0.05 mol dm-3, to ~CTAC Micelles, [C1-] = 4.00 rnol dm-Following Moller and Kragh-Han~en~~the differences between pK," (or pk") in CTAC micelles (4.00 mol dm -KCl) and in CPC micelles (0.05 rnol dm-3 KCl) are denoted as AApKf: (or AApk"). The AApKf: C-t(O.08-O.l8)] vary from 1.11 (AApKr, of ethyl eosin) to 2.94 (AApKf:, of eosin).As follows from eqn. (11)-(13) both pk" and log a cause the resulting effect. [Note that pKf:, values of eosin and ethyl eosin in CPC micelles (Table 1) are expressed in the pH, scale.] A special case for fluorescein (AApKf:, = -0.38) can be elucidated by the fact that in CPC micelles (0.05 mol dm-3 KCl) only the H2R (not H3R+) species are solubilized. As a whole the AApk" values vary within a narrower range than Apk" (Table 4); lowering the potential of the Stern layer leads to ca. 10-1000-fold decrease in the k" values. The effect differs for phenolic and carboxylic groups: (i) The average value for eight AApk$H is 1.4 & 0.3 (the value AAPG,~~is excluded, see above).For eight sulfonephthaleins' the AApk;,.. values are in the range 1.63-1.94; (ii) The average AApk:o2H value is 2.3 f0.5 (four pk" values). Thus if the decrease in Stern layer potential is evaluated as A'P = 59AApk" (27) rather close values (96 to 114 mV, average: 107 f7 mV) may be obtained for eight sulfonephthaleins, but if using AApk;, C02H (fluorescein) and AApk;, COzH (2,7-dichloro-fluorescein, eosin) the values AY = 142 and 130-171 mV differ more significantly (the calculations are conventional as the pk" values in CPC micelles are obtained at 20°C). The observed effects can be explained through the peculiarities of the solubilization of carboxylates in cetylpyridinium micelles (the similarity of pKr in CTAC and CPC micelles at equal [Cl-1, values is still proved only for ~ulfonephthaleins~~), J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 or/and it must be assumed that the rise in [Cl-J, results not only in Y lowering. The salt additives are known sometimes to increase the thickness of the Stern layer.' Besides, with increase in the counter-ion bulk concentration, the micelles growth and their changes in shape take place to some extent in chloride systems too.2 Note that the growth of micelles occurring both during c, increase (DTAB) and addition of salts to DTAC and DTAB (4 mol dm-3 NaCl and NaBr) leads to the decrease in Eq values;" for CPB and CTAB micelles (c, = 0.05 mol dmP3) without salt additives the E: values practi- cally coincide." Thus it can be supposed that the properties of the Stern region at the increase in ionic strength of the bulk phase became still more 'non-aqueous'.This may be the reason for the additional increase of pk;,,, as compared to p@, (see the AApk" in Table 4). The K, values of all the three dyes increase by ca. three times their original values in transferring from one micellar system to the other (Table 5). Moreover, the decrease of pK,, -pK,, values in the case of fluorescein and their increase in the case of eosin (in both cases with reference to aqueous medium) are more marked in micellar solutions with high [Cl-1, values [Table 5; see also eqn. (25) and (26)]. Following this logic the pKi (and pk') values at high ionic strength of the bulk phase may, generally speaking, differ from those at low electrolyte concentration.Further research is necessary in order to clarify this problem. Organic anions inside the micelles are supposed to be associated with cationic head gro~ps.'~'~' In this case the competition between CO; groups (with localized charge) and C1- may occur in their interactions with cationic surfactants. If the increase in [Cl -1, results in 'dye-surfactant associates' dissociation as well (as a consequence of further saturation of the Stern layer with the ele~trolyte'.~), then the contribution of the formation of such associates to the medium effects (ApKf, Apk") at low ionic strength of the bulk phase should be recognized. Such additional stabilization of carboxylates can be an additional source of hindrance in attempts to model pKi and pKY using pK, in water-organic compound mixtures.The errors (sometimes k0.2) and number of AApk" obtained (Table 4) do not allow more definite conclusions to be made. On the other hand if the CO, and SO, groups of HR-species (5a, 9) can be regarded as 'neutralized' with surfactant cations, then the pK;, values of fluorescein and sul-fonefluorescein must be equal to the pKr, value of 6-hydroxy-9-phenylfluoron (Scheme 4). But in CPC micellar solution (0.05 mol dm-3 KCl) this pK;, is found to be 4.67 & 0.07,30-32 i.e. 0.8-0.9 pK, units lower than pK;, of fluorescein and sulfonefluorescein. In water this difference, 6 [see eqn. (15)], is 0.5.19It can be considered as evidence of oversimplicity of the model of complete neutralization of anionic groups in the Stern layer of cationic surfactants.Conclusions According to the electrostatic approach, reflected in eqn. (l), increase of the counter-ion concentration in the bulk phase to 13 14 Scheme 4 Ionization of 6-hydroxy-9-phenylfluoron 4 mol dm-3 leads principally to micellar charge shielding and to the reduction of the contribution of the quantity YF/ 2.3RT. It allows the consideration of the set of pK: and k" values as a rough approximation to that of pKi (pk'). In several cases the values are essentially higher than pKz (pk"). The influence of salt additives on the specific interactions cannot be entirely excluded either.The Stern region is known to be highly aqueous, contain- ing high concentrations of ionic head groups and counter ions.1~2*6~10*3'*5'However, the equilibrium parameters, deter- mined for the solubilized oxyxanthene series, differ signifi- cantly from those in water or in water-salt systems. Three main features, typical for organic microenvironments, can be pointed out: (i) red shift of absorption and emission bands for species of type 7;(ii) differentiating effect of the pseudophase; different medium effects, Apk, reflect the charge type of the acid-base couple and the nature of ionizing group (ApPoH < Apkco2"); (iii)shifts of tautomeric equilibria. All these factors, controlling the intrinsic pKi values, cannot be caused by anything except the proximity of the hydrocar- bon core.It is difficult to compare directly the pk" (~pk') values with those in water-organic compound mixtures because of the extrathermodynamic nature of yH values. However, the data give rise to some doubts about the correspondence of the pseudophase solvent properties to those of a certain water- organic compound mixture. The pk;, and pk;,,, values of solubilized dyes show some- what differing increases at increasing ionic strengths of the bulk phase. A preliminary hypothesis for explanation of this and some other effects is that the micellar pseudophase (at least the Stern region) becomes more 'non-aqueous' at high salt concentrations in the water phase. Note Added in Proof In the paper of N.0. Mchedlov-Petrossyan and R. Salinas Moyorga, J. Chem. Soc., Faraday Trans., 1990, 90, 3025, the following corrections should be made: Within Scheme 1, "97 \ Aorokshould appear as References 1 E. J. R. Sudholter, G. B. Van de Langkruis and J. B. F. N. Eng-berts, Red. Trav. Chim. Pays-Bas, 1980, 99, 73. 2 C. A. Bunton and G. Savelli, Adv. Phys. Org. Chem., 1986, 22, 213. 3 P. Mukerjee and K. Banejee, J. Phys. Chern., 1964,68,3567. 4 N. Funasaki, Nippon Kagaku Kaishi, 1976, 722. 5 F. Dorion, G. Charbit and R. Gaboriaud, J. CoZloid Interface Sci., 1984, 101, 27. 6 L. S. Romsted, J. Phys. Chem., 1985,89, 5107. 7 G.S. Hartley and J. W. Roe, Trans. Faraday SOC.,1940 36,101. 8 M. S. Fernandez and P. Fromherz, J. Phys. Chem., 1977, 81, 1755.9 M. E. Diaz Garcia and A. Sanz-Medel, Talanta, 1986,33, 255. 10 N. Funasaki, J. Phys. Chem., 1979,83, 1998. 11 G. Charbit, F. Dorion and R. Gaboriaud, J. Chim. Phys., 1984, 81, 187. 12 C. J. Drummond, F. Grieser and T. W. Healy, J. Chem. SOC., Faraday Trans. 1, 1989,85,521. 13 C. J. Drummond, F. 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ISSN:0956-5000    

DOI:10.1039/FT9949000629

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(4), 641-647 Computer Modelling of Phosphate Biominerals Transfer of Parameters for Interatomic Potentials for Different Polymorphs of Divalent Metal Diphosphates Marina G. Taylor and Kenneth Simkiss Department of Pure and Applied Zoology, University of Reading, P.O. Box 228,Whiteknights, Reading, UK RG6 2AJ Maurice Leslie SERC Daresbury Laboratory, Daresbury, Warrington, UK WA4 4AD Manganese, zinc and calcium diphosphates have been modelled using atomistic simulation techniques. The ionic model was used for the cation-anion interactions. Parameters for the short-range interatomic potentials previously calculated using the electron gas methods for the diphosphate anion of or-magnesium diphosphate, were used together with specific metal oxygen parameters.The structural data for each compound were repro- duced well, indicating that the anion parameters were transferable. These studies present new opportunities for modelling important aspects of biominerals which are not readily accessible to experimental studies. Calcium phosphates form the major inorganic component of mineralised tissue such as bone and teeth in vertebrates. Calcium phosphate appears as hydroxyapatite, frequently with some carbonate included in the lattice. Precursors to the final mineral have been proposed ranging from an amorp- hous calcium phosphate phase, found in matrix vesicles close to the mineralising front, through to brushite and octacal- cium phosphate. In addition, in many invertebrates tissues intracellular granules, which also have an amorphous struc- ture, are found which have an ionic composition which can be an orthophosphate or a condensed phosphate such as diphosphate (pyrophosphate).' The cations present are mostly calcium and magnesium but a feature of these bio- minerals is their ability to incorporate other cations such as manganese, cobalt, zinc and transuranics when present in the environment and so these materials can be seen as pro- viding a sink for pollutants and so detoxifying them.In the bone precursor, in which magnesium and zinc have been detected in addition to calcium,' the magnesium may have a role as a crystal inhibitor. The amorphous materials are pro- duced from aqueous solutions at environmental temperatures but have many of the properties associated with glasses. They are of considerable biological importance since they possess unusual solubility and solid-phase properties that have been exploited by living systems in a variety of mineralising, acid- base regulating and detoxifying systems.Analytical studies have established the composition of the materials and provided information on the nature of the phosphate species. Structural studies using X-ray absorption spectroscopy have been used to determine the local atomic structure around calcium and other cations in granules which have orthophosphate and diphosphate anions. 1*3-5 The dis- tribution of cations in these materials is not known but may be similar to those found in mixed alkali-metal-silicate Atomistic simulation methods of compounds with known crystal structures are being used to determine the sub- stitution energies of the dopant metals and subsequently the surface reactivity of these biominerals.These methods have been used successfully to model the static and dynamic properties of the perfect lattices of compounds such as zeo- litic silica polymorphsg and calcium carbonate." We have previously modelled the perfect lattice of cr-magnesium diphosphate" and we now present our results on the trans- ferability of parameters, determined for this magnesium com- pound, to other anhydrous diphosphate structures with calcium, manganese and zinc as the counterions. Structures Inorganic diphosphates exhibit two major classes of polymorphism" with differing structural features.Diphos- phates are condensed dimers of phosphate units with a formula of P,O$- which consists of two PO, units linked by an oxygen bridge in the P-0-P form. The conformations of the PO, units, staggered or eclipsed and the angle of the P-0-P bridge are the main features which characterise the two major types, thortveitite and dichromate structures. The thortveitite structure has the staggered conformations of the PO, groups and the P-0-P bridging angle greater than 140". Only in the dichromate-type structure does the bridging oxygen come into the first coordination sphere of the cation. When the structure of thortveitite, Sc,Si,O,, was first deter- mined, it was reported that the Si-0-Si bond angle was linear, i.e.180°.'2 This has been the subject of much dis- cussion, especially as other similar compounds have mostly non-linear bridges, but the linear bridge has subsequently been confirmed." Among the diphosphates several linear P-0-P bridges have been reported'3,'4 but it is generally considered that this might arise as a result of thermal dis- order. We report the modelling of various metal diphosphate compounds. The initial study was of a-magnesium diphos- phate, a compound crystallising with a thortveitite type struc- ture.' ' These studies have been extended to manganese and zinc diphosphates which are isomorphous with magnesium diphosphate and to /3-calcium diphosphate which is not and has the dichromate structure.The study was extended, in part to test the transferability of the parameters for interatomic potentials, but also because these compounds are relevant to our experimental studies of biominerals, particularly those related to the incorporation of foreign cations into the calcium magnesium diphosphate granules. Manganese diphosphate is unique in being the only diphosphate with the thortveitite structure at room tem-perature. In early X-ray studies of manganese diphosphate, a linear P-0-P bridging bond was indicated.', A more recent investigation of the structure' has suggested that there may be thermal disorder in the bridge and a split-atom model was proposed retaining the C,,, symmetry but with a bridging angle of 165.9'.Only an occupancy factor of 0.80 was found and even further disorder was suggested for the remaining fifth. A further structure with a bent model with the symmetry reduced to C, and a bridging angle of 164.5" could not be refined in the X-ray study. We have attempted to model the structures with a linear P-0-P bridge and the split atom model. The second set of diphosphates we have considered are the a-and p-zinc diphosphates which also crystallise in thortveitite type structures. Both forms are pre- pared from a melt obtained from the decomposition of zinc ammonium phosphate. For the higher temperature form, p-zinc diphosphate, a linear bridging P-0-P has also been proposed with the suggestion that this too may reflect thermal disorder though in this case there have been no further attempts to refine the crystal structure taking this into account.l4 The lower temperature phase, a-zinc diphosphate is described as a hexpartite structure in which the a-axis is tripled and the c-axis is doubled relative to the thortveitite structure.16 It is believed to have the most complex structure of the series of metal diphosphates. The structure consists of a sequence of layers, one of which is like that found in a- copper diphosphate followed by two layers like those found in a-magnesium diphosphate. In a-magnesium diphosphate both the a-and c-axes are doubled compared with the man- ganese structure.' 5,1 Calcium diphosphates have very complicated structures with coordination numbers ranging from seven to nine, calcium oxygen distances ranging from 2.250 to 2.807 8, in two eight-coordinate sites in the high-temperature ct-form'* and 2.318 to 2.927 in four different coordination sites with seven, nine, seven and eight nearest-neighbour oxygens in the lower temperature p-form.lg In the latter case in sites one and four, bridging oxygens come within 3 8, of the central calcium.The a-form has the thortveitite-type structure and the p-form the dichromate type structure. All the crystalline calcium diphosphates are high-temperature phases. The p-calcium diphosphate was model- led partly because it is the lower temperature form, but also becuse it crystallises in the dichromate structure, and is tetragonal whereas the magnesium, manganese and zinc crys- tals are of the thortveitite type and are monoclinic.In the calcium structure there are two types of diphosphate units each with an eclipsed configuration of the PO, units. The structure shows infinite chains of diphosphate units linked in all directions by calcium ions with calcium found in four dis- tinct sites. Methods Computer Modelling Computer modelling was carried out using the University of Reading's AMDAHL 5870 and a SUN SPARC station 1. The programs used were THBREL and THBPHON made available by the SERC CCP5 scheme. These programs were developed initially at AEA Harwell to study defects in solids20*2' using the Mott-Littleton approximation.22 Theoretical Methods The compounds were modelled as perfect ionic lattices with cations interacting with internally covalent bound diphos- phate anions. Two- and three-body terms were included to model the phosphorus-oxygen bonds in the anion.The lattice energy can be computed at constant volume and constant pressure as the sum of the long-range Coulom- bic interactions, the short-range non-bonded potentials and the force field terms for the covalent anion. = ECoulornbic + Eshort-range + Etwo +three body The lattice energy is minimised using classical optimisation techniques23 such as the Newton-Raphson method as a func- tion of structural parameters, lattice parameters and atomic J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 coordinates.At the minimum the derivatives of the energy with respect to the geometric factors will be zero. Bond lengths and bond angles are calculated and compared with the crystallographic data to give an indication of the good- ness of the potentials used. A further check on the inter- atomic potentials is obtained by generating the phonon spectrum. The eigenvalues calculated from the THBPHON program using a wave vector of 0.0o0, O.OO0 and 0.001 have been compared with the infrared spectra of metal diphos- phate corn pound^.^^ The starting parameters for the non-bonded interatomic potentials were mean values obtained by the electron gas method as previously de~cribed.~~-~' Values were calculated for each atomic pair in magnesium, manganese, zinc and calcium diphosphates.When required, more than one range of parameters was used for different cation-oxygen distances. Initially the charges and parameters which were determined earlier by the electron gas technique for the diphosphate anion in magnesium diphosphate have been used' and have been maintained as far as possible throughout. Some changes were made to the charge distribution within the anion to improve the fit in some models. Specific values have been cal- culated for the parameters for each cation-anion non-bonded repulsion contribution to the interatomic potential of the compounds of interest. For the electron gas potential calcu- lations the wavefunctions were calculated in the Madelung potential well to account for the local crystal environment.The method used for magnesium diphosphate' ' was to set up different charge models for phosphorus and the bridging and terminal oxygens in each crystal and calculate the Madelung potential. The electron densities were then determined from expansions of Slater functions modified to account for the Madelung potential experienced in the crystal. The charge on the cations was maintained as 2+. The charges on the atoms in the anions were initially those which were used to model magnesium diphosphate. These were the formal charges in the Langmuir sense, i.e. 1+ on phosphorus, 1-on the ter- minal oxygen and 0 on the bridging oxygen. These were adjusted within the anion to improve the fit obtained from the calculation when necessary.The parameters for the non- bonded repulsive potentials in the anion were maintained throughout, except where indicated in Table 1. Clementi free- ion wavefunctions were used for manganese and calcium as had been used previously for magnesium. There were no Table 1 Parameters for the Buckingham non-bonded interatomic potentials for static lattice simulations of diphosphate compounds species AIeV pIA C/eV A6 range/A Mg2+-0, 3551.55 0.243 32 6.00 0-20 Mn2+ -0, 2538.565 0.267 56 4.00 0-2.2 3028.565 0.267 56 4.00 2.2-20.0 a-Zn2+ -0, 1052.000 0.287 82 0.00 0-20 P-Zn2+ -0, 1107.000 0.287 82 0.00 0-2.15 1125.000 0.287 82 0.00 2.15-20 P-Ca2+ -0, 3000.750 0.272 98 0.00 0-2.40 3100.833 0.272 98 0.00 2.40-2.60 3500.750 0.272 98 0.00 2.60-20.0 fi-Ca2+ -0, 500.53 1 0.276 00 30.00 0-3.10 750.531 0.276 00 00.00 3.10-20.0 OT-0, OT-0, OFl--OB 1394.529 108 1.8 17 764.00 0.297 82 0.303 57 0.320 00 120.00 80.00 53.00 0-20.0 0-20.0 0-20.0 P-P 1030.43 0.356 62 0.00 0-20.0 P-0, 388.475 0.399 78 2.00" 0-20.0 P-0, 300.7 16' 0.428 76 2.00* 0-20.0 a.b Values of 10.00 and 0.00, respectively, were used for the calcium compound.A value of 370.716 was used in the manganese com- pound. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 643 7.5 7.5 6.5 6.5 5.5 5.5 4.5 4.5 23 3.5 233.5 5: 5: 2.5 2.5 1.5. 1.5 0.5. 0.5 -0.5. I 1 I 2 I 3 1 4 -0.5 0 1 2 3 4 r/A r/A 7.5 7.5- 6.5 6.5 - 5.5 5.5 - 4.5 4.5 - % i.3.5r -..>,.. T 3.5-5 2.5- 2.5 - 1.5 1.5- 0.5 0.5 - -0.5. I 1 I 2 I 3 I 4 - 0 0 . 1 5 2 . 1 3 , 4 r/A r/A Fig. 1 Mg,Mn and Zn overlap except where indicated. V(r)for non-bonded repulsive terms us. interatomic distance for (a)0,-0,, (b)0,-0,, (c) P-0, and (d)P-0, . The curves for Ca, parameters for Zn2 + in Clementi-Roetti28 so the cation density function was calculated n~merically~~ with a rela-tivistic correction. In the electron gas method exchange energy terms were calculated using the Handler formula and the correlation energy terms were calculated using Wigner’s function. The electron gas data were then fitted to the analy- tic Born-Mayer function.The electron gas method produces a range of parameters depending on the local environment of each individual atomic species. The starting point for fitting the structures used a mean value of the calculated values of the A and p parameters. Any modifications to the A parameter during the fitting were, as far as possible, kept within the range of the calculated values. Plots of the Born-Mayer function using the parameters for the anion in each compound for OT-OT, OT-O,, P-0, and P-0, interactions are com- pared in Fig. 1. 0, represents a terminal oxygen and 0, a bridging oxygen. Although the numerical values of A and p were not the same for each compound, these parameters are correlated and it is seen that the potential energy is virtually identical for O,-OT and 0,-0, interactions and there are only small differences for the P-0, and P-0, interactions. Therefore these parameters obtained for magnesium were used in modelling the anion in the other compounds.Van der Waals terms for the short-range dispersion effects were calculated using the formula derived by Slater and Kirk- For the interaction between species i and j C,(ij) = 3/2ai aj/[(ai/Pi)”2 + (c~~/P~)”~] Table 2 Force constants used in the perfect lattice simulations species Mgz+ Mn2+ a-Zn2+ b-ZnZ+ B-CaZf two-body P-0, k/eV k2 25.0 25.0 25.0 25.0 25.0 bond length, rlA 1.590 1.568 1.589 1.569 1.615 P-0, kleV A -35.0 30.0 35.0 30.0 35.0 bond length, r/A 1.516 1.526 1.52 1 1.555 1.518 fraction of Coulombic term excluded 0.35 1.0 0.45 0.50 0.50 three body 0,-P-0, k/eV rad-2 10.0 15.0 12.5 12.0 10.0 mean equilibrium angle, eldegrees 112.33 112.50 1 12.09 111.75 1 12.60 0,-P-0, k/eV rad-2 15.0 16.0 12.5 16.0 15.0 mean equilibrium angle, eldegrees 106.33 106.2 106.65 106.50 105.4 P-0B-P k/eV rad-2 20.0 20.0 20.0 20.0 25.0 mean equilibrium angle, eldegrees 144.0 180.0 143.0 180.0 134.2 where ai is the static polarizability of species i and Pi the effective number of electrons contributing to the polarizabil- ity.A damping factor zDi,(r) was also included to reduce the dispersion energies when the wavefunction overlap is not neg- ligible.31 The short-range terms can be collected into the Buck- ingham potential The long-range electrostatic interactions between each ion pair are included with an additional term 4i4j/rij where 4i represents the charge on ion i and rij is the interionic dis- tance.An approximate calculation of the two body force con- stants was made by considering the energy of assigned vibra- tions in the infrared spectrum. The P-0 bonds were treated as harmonic oscillators Fj = ik(r -ro)2 Table 3 Charge models for diphosphates species Mg Mn or-Zn B-Zn /?-Ca cation 2+ 2+ 2+ 2+ 2+ phosphorus 1+ 1,8+ 1.05+ 1.80+ 1.2+ bridging" 0.0 -0.10 -0.10 -0.10 -0.40 oxygen, 0,terminal -1.0 -1.25 -1.00 -1.25 -1.00 oxygen, 0, " THBPHON program requires a non-zero charge on each species. J. CHEM. SOC. FARADAY TRANS., 1994, VOL.90 Table 4 Perfect lattice simulation of manganese diphosphate (Mn,P,O,); comparison of experimental (X-ray) and calculated values (a) Cell parameters and lattice energy cell parameters calculated experimental alA 6.42 6.63 blA 8.59 8.58 CIA 4.42 4.55 /?/degrees 103.2 102.7 lattice energy/eV -88 x 2 -87 x 2" (Z = 2) (b) Bond distanceslA and anglesldegrees species X-ray CONV CONP Mn-0, X 2 2.164 2.149 2.148 Mn-0, x 2 2.132 2.147 2.134 Mn-0, x 2 mean value 2.316 2.204 k0.080 2.313 2.203 f0.078 2.315 2.199 f0.082 P-0, 1.568 1.564 1.561 P-P 3.140 3.128 3.121 P-0, 1.519 k0.001 1.519 1.516 k0.005 <P-0,-P 180.0 180.0 180.0 <OB-P-o+ 106.2 & 1.3 107.2 f0.2 102.3 & 0.5 mean value <o,-P-0, 112.3 f0.3 11 1.0 f0.8 11 1.5 f0.1 mean value " From Born-Haber calculation.where (r -r,) represents the displacement from the equi- librium distance ro/A and k is the spring constant. The three- body force constants were then estimated by considering the assigned vibrations in relation to the two-body terms. where k, is the bond bending force constant and 8, is the equilibrium bond angle. Results The parameters for the non-bonded terms of interatomic potentials used for magnesium, calcium, manganese and zinc diphosphates are listed in Table 1. The force constants for the two- and three-body terms are also given (Table 2). The geometric fitting of manganese diphosphate with the linear bridge was good.The crystal structure shows man- ganese in octahedral coordination with two Mn-0 distances at 2.132 8, and two at 2.164 A. The other two bond distances were longer at 2.316 A and so two ranges of A parameters were required for fitting the Mn-0 bond lengths. These were selected on either side of the calculated mean value. The distribution of charge within the diphosphate anion was modified from the values used for the a-magnesium diphos- phate to improve the fit but the overall charge of the anion was maintained (Table 3). Calculated bond lengths in the structure relaxed at constant volume and pressure are given in Table 4 and compared with those determined by X-ray crystallography and neutron diffraction studies. There was a decrease in the cell dimensions of the relaxed structure com- pared with the X-ray determinations (Table 4).The calculated lattice energy, -88 eV compares with the value determined J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 5 Perfect lattice simulation of a-zinc diphosphate (Q-Zn2P20,); comparison of experimental (X-rays) and calculated values (a) Cell parameters cell parameters calculated experiment a1 a/A 19.67 20.07 blA 7.86 8.26 CIA 8.63 9.09 flldegrees 105.8 106.4 lattice energy/eV -79.7 x 6 -71.0 x 6" (2= 6) (b) Bond distance/A and anglesldegrees ~~ species X-ray CONV ~~ CONP Zn(Cu)-0, 1.956 1.955 1.956 2.024 2.053 2.034 2.060 2.057 2.040 2.060 2.065 2.063 2.089 2.117 2.090 mean value 2.038 f0.046 2.049 f0.052 2.037 f0.045 Zn(Mgl)-oT 2.032 2.05 1 2.025 2.056 2.022 2.044 2.083 2.104 2.089 2.097 2.108 2.104 2.102 2.111 2.107 2.177 2.132 2.1 14 2.090 f0.046 2.089 f0.037 2.080 f0.035 1.928 1.951 1.945 2.025 2.002 1.998 2.032 2.087 2.042 2.056 2.115 2.105 2.093 2.115 2.117 2.027 f0.055 2.054 f0.066 2.041 _+ 0.065 1.603 1.578 1.595 1.599 1.586 1.595 1.566 1.580 1.596 3.Ooo 3.012 3.045 3.046 3.022 3.049 1.529 & 0.012 1.519 f0.004 1.511 f0.004 1.512 f0.016 1.519 f0.006 1.511 f0.005 1.523 f0.006 1.527 f0.003 1.516 f0.005 Plopl 139 145 145 p2,3°p2.3 148 145 146 111.8 f1.5 113.1 f3.5 109.0 f 1.0 106.9 & 3.8 109.1 f2.5 109.9 f1.6 111.2 f2.0 109.9 f1.3 109.2 f0.6 106.2 k 3.5 109.0 f2.9 109.7 f1.7 112.0 f0.4 109.6 k 1.9 108.7 f0.5 106.8 f2.3 109.3 f2.2 110.2 f2.2 a From Born-Haber calculation. from experimental data in a Born-Haber type calculation -87 eV. Although there are no literature values for elastic constants and relative permittivities the values obtained were judged reasonable.We were unable to fit the split-atom model containing the disordered bridging oxygen with THBREL. The structure relaxed but a new site for man-ganese binding appeared which was inconsistent with the crystallographic structure proposed by Stefandis and Nord.' Both a- and /.?-forms of zinc diphosphate were fitted. The charge distribution in the diphosphate anion in the a-form was maintained as in magnesium diphosphate.The param- eters for the non-bonded interatomic potentials in the anion were those used for modelling magnesium diphosphate. Some Table 6 Perfect lattice simulation of ,!&zinc diphosphate (p-Zn,P,O,); comparison of experimental (X-ray) and calculated values (a) Cell parameters and lattice energy cell parameters calculated experimental 4 6.57 6.6 1 biA 8.15 8.29 CI'A 4.39 4.5 1 Ptdegrees 105.1 105.4 lattice energy/eV -77 x 2 -71 x 2" (2= 2) (b) Bond distances/A and bond anglesldegrees species X-ray CONV CONP ~ Zn-0, x 2 2.00 1 2.001 1.998 Zn-0, x 2 2.06 1 2.067 2.080 Zn-0, x 2 2.275 2.245 2.27 1 mean values 2.112 & 0.1 18 2.104 f0.103 2.1 16 & 0.1 14 P-0, 1.569 1.546 1.556 P- P 3.137 3.093 3.1 13 mean value P-0, 1.555 f0.001 1.574 f0.010 1.569 f0.008 angles POBP 180.0 180.0 180.0 mean value %PO, 111.5 0.8 107.4 f3.7 108.7 & 1.9 110.2 f2.0 108.3 f1.2 110.6 f1.4 " From Born-Haber calculation.small modifications were made to the force constants for the two-and three-body terms (Table 2) to improve the fit obtained. The parameters used initially for the zinc-oxygen interatomic potentials were those calculated using the elec- tron gas method. A lower A value, however, improved the fit. In the a-form16 zinc is in three different sites with coordi- nation numbers of five, six and five. This feature was repro- duced in the relaxed structure.The zinc-oxygen and phosphorus-oxygen distances are given in Table 5 and are in reasonable agreement with the crystallographic data. The lattice energy and cell dimensions of the relaxed structure are compared with the experimental data in Table 5. Once again reasonable elastic constants and relative permittivities were obtained although there is no reference data available. The crystallographic struct~re'~ of the p-zinc diphosphate was solved with a linear P-0-P bond although the ques- tion of thermal disorder was discussed. As with the structure of manganese diphosphate, also with a linear bridging oxygen, the structure was successfully modelled only when the charges in the diphosphate anion were similarly modified. Two ranges of the parameters for the zinc-oxygen non-bonded terms were also required to reproduce the crystallo- graphic geometry (Tables 1 and 6).The calcium diphosphate structure did not model as well as magnesium, manganese and zinc diphosphates. The charge model was changed slightly so that the bridging oxygen carried a partial charge of -0.4 and the charge on the phos- phorus atoms was then increased to + 1.2 to maintain charge neutrality. The A parameters are consequently from the higher range of calculated values. The bond distances for the four calcium sites and the phosphorus-oxygen distances are given in Table 7. Most of the values compare well with X-ray crystallography, especially the nearest neighbours. There was a tendency for distant terminal oxygen atoms in sites one and three to move in closer to calcium and increase the coordi- nation number.Despite the proximity of the bridging Table 7 Perfect lattice simulation of p-calcium diphosphate (p-Ca,P,O,); comparison of experimental (X-ray) and calculated values (a) Cell parameters and lattice energy Cell parameters calculated experimental ~ a/A 6.590 6.684 6.590 6.684blA 24.144 24.144CIA (a = B = y)/degrees 90 90 lattice energy/eV (2 = 8) -73.8 x 8 -76.5 x 8" (b) Bond distances/A and bond anglesldegrees species X-ray CONV CONP Ca(1)-0, 2.340 2.364 2.332 2.360 2.315 2.334 2.369 2.365 2.336 2.409 2.380 2.362 2.416 2.455 2.449 2.457 2.556 2.512 mean value 2.392 & 0.039 2.442 f0.105 2.385 f0.021 2.814 2.89 1 Ca(l)-oB Ca(2)-0, 2.780 2.342 2.397 2.364 2.42 1 2.351 2.378 2.414 2.455 2.469 2.509 2.470 2.518 2.557 2.55 1 2.539 2.640 2.569 2.565 2.668 2.649 2.666 2.745 2.661 2.667 mean value 2.534 f0.135 2.519 f0.109 2.855 2.680 mean value 2.570 k0.162 2.535 f0.107 Ca(3)-0, 2.318 2.330 2.332 2.341 2.309 2.332 2.343 2.383 2.352 2.356 2.469 2.477 2.462 2.57 1 2.501 2.539 2.670 2.636 2.692 2.747 2.683 mean value 2.434 f0.129 2.502 k 0.153 2.470 f0.138 2.760 2.727 ca(4)-0T 2.370 2.371 2.342 2.372 2.333 2.345 2.397 2.386 2.383 2.435 2.43 1 2.405 2.466 2.478 2.45 1 2.510 2.480 2.477 2.794 2.48 1 2.493 mean value 2.477 k0.138 2.424 f0.054 2.412 f0.058 ca(4)-0B P1-0, P3-0, '2-OE p4-0BP,--Pi 2.927 1.637 1.617 1.590 1.616 2.955 1.607 1.61 1 1.612 1.610 2.98 1 1.603 1.610 1.611 1.608 2.973 p3-4mean values 2.991 2.992 2.986 coplanar bonds 4 P-0, 1.497 f0.011 1.495 f0.007 1.491 f0.002 8 P-0, 1.529 f0.016 1.500 f0.003 1.497 k 0.002 angles ploEp2 p3°Bp4mean values 130.5 137.8 135.8 136.5 135.4 136.1 OJ'iOT OTPZ0T 112.6 f3.2 111.8 k 1.3 110.2 f2.5 109.7 k0.6 110.2 f2.7 109.6 k0.2 OTP3OT 111.9 & 2.5 109.6 ,+ 0.9 109.5 & 0.4 OTP4OT OBP2OT OBP3OT OBPiOT OBP~OT 113.9 f 1.2 106.1 & 3.2 106.0 k 1.5 106.8 f 1.5 104.7 k2.2 110.3 k 3.5 108.7 1.2 109.3 & 0.4 109.5 f0.9 108.6 & 1.3 110.2 f 3.7 108.7 f0.9 109.4 k0.3 109.4 f0.4 108.6 f1.3 a From Born-Haber calculation.J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 8 Highest frequencies calculated from the THBPHON program compared with the highest experimental frequency band in the infrared spectrum assigned to v(P-0,) compound calculated frequency/cm- experimental infrared frequencylcm - 1210 1200 1230 1230 1212 oxygens to the central calcium in two sites as determined by X-ray crystallography our model did not allow such a close approach. Most of the lattice parameters decreased slightly on relaxation. The modelled structures were all stable in that there were no imaginary eigenvalues in the THBPHON calculations with a wave vector of q = O.OO0, O.OO0 and 0.001. The frequencies (wavenumbers) derived from the eigenvalues were in the same range as found in the experimental infrared spectra of all the compounds and compared with the range of values calcu- lated in normal coordinate analyses of divalent metal diphos- phate~.~~Because of the limitations of computation at that time these analyses used a model with D,, symmetry despite the fact that the symmetry of the diphosphate anions with a bent bridge can be as low as C,.For these reasons a more detailed analysis does not seem to be justified. The highest frequencies calculated from THBPHON relating to the phos- phorus terminal oxygen stretching vibrations do, however, compare reasonably with experimental values in Table 8. Discussion We have presented our results of modelling different metal diphosphates using the parameters for the non-bonded inter- atomic potentials obtained using the electron gas method- ology and fitted to the Born-Mayer analytic function and dispersion terms calculated by the Slater-Kirkwood method.We have previously used the parameters for the diphosphate unit in our modelling of magnesium diphosphate and have now tested their transferability to model other similar com- pounds. They were used most successfully in modelling the compounds which are structurally isomorphous with magne- sium diphosphate, the manganese and zinc diphosphates. The linear P-0-P bond angle was modelled successfully in the manganese and /%zinc structures. The P-0-P angle has been discussed in the context of the dr-pn overlap and results in the shorter P-0, bonds32 which are found in the manganese and zinc structures with the linear P-0-P bond.The linear bridging oxygen atom did present a slight problem as it is believed that this is an average position hiding the thermal disorder. However, an attempt to model the disordered structure, the split-atom model of manganese diphosphate proposed in an X-ray study was not successful. A double cell was set up to accommodate the extremes of the disorder but this resulted in the appearance of an additional site for manganese in our relaxed structure. Only in the model with a linear bridge was one unique MnO, polyhedron found. The structures with a linear P-0,-P bond in the manganese and the p-zinc compounds needed a modification in the charge distribution in the anion and two ranges for the metal oxygen parameters to reproduce the structure satisfac- torily. The rather more complicated a-zinc diphosphate mod- elled extraordinarily well with the same charge distribution and very minor changes in the parameters used in the magne- sium study.Calcium has a bigger ionic radius than magne- J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 sium, manganese or zinc and so is able to accommodate a higher number of nearest neighbours. In the X-ray structure of /?-calcium diphosphate a judgement was made as to what constituted the nearest neighbours of calcium. On the basis of a distance of under 3 A it was found that the coordination numbers were seven, nine, seven and eight for the four inde- pendent calcium sites.Included in the coordination sphere of Ca (1) and Ca (4) was a bridging oxygen atom at 2.78 and 2.93 A, respectively. All the sites are described as being dis- torted. Sites one and three are described as distorted pentag- onal bipyramids or an octahedron distorted by the approach of a more distant oxygen. Site three is described as a pentag- onal bipyramid with the apical atom replaced by three atoms and the fourth site as an octahedron with two oxygen atoms jammed into edges. Despite this challenging structure the model was reasonably successful especially at the closest neighbours. Conclusions Several metal diphosphates have been modelled using the same parameters for the non-bonded interatomic potentials in the anion as previously used to model magnesium diphos- phate.Phosphorus-oxygen bond distances tended to be slightly short but most were within 2% of the X-ray determi- nations. In the calcium structure the range of phosphorus ter- minal oxygen bond lengths was 1.480 to 1.562 A. Only the unusually long bond distance was not reproduced accurately. Bond angles in the diphosphate units tended to move towards the standard tetrahedral value of 109". The metal- oxygen distances, in general, were in good agreement with the crystallographic data. The bond distances in the manganese and /%zinc diphosphate structures relaxed at constant pres- sure were within 1.35% of the crystallographic values while those in a-zinc diphosphates, which had a more complicated structure, had a difference of 2.89% in one bond.Even in the calcium diphosphate structure most bond distances were within 2% of the values determined by X-ray crystallography. The work has demonstrated that the interatomic potentials used for the static simulation of the perfect lattice of diphos- phates can be used with confidence in various metal com- pounds provided specific potentials are used for the metal- oxygen interactions. Changes in the cell dimensions show that there are decreases in cell volumes on relaxation at constant pressure compared with the experimental values. Most of the experi- mental lattice parameters were obtained at ambient tem-perature except for B-zinc diphosphate when measurements were made on a heated cry~tal'~ (temperature not reported).The potentials have been derived from the electron gas meth- odology using expansions of Slater functions, where there is no compensation for the effects of temperature as there would be for empirically fitted parameters using the CCP5 program THBFIT which maintains cell volumes. Improvements may be obtained by using the shell model33 which takes account of the polarizability of the ions, notably in this case the oxygen which has a more diffuse electron cloud. These studies indicate that these potentials can be used to investigate cation substitutions in different diphosphate compounds and so extend our studies to model biological systems.The potentials are also being evaluated for ortho- phosphate compounds so that we may then model the struc- ture and reactivity of surfaces of the many important phosphatic biominerals such as hydroxyapatite. We acknowledge the help of Dr. A. H. Harker, AEA Tech-nology, Harwell in the calculation of interatomic potentials using the electron gas models. We thank Drs. M. G. B. Drew and P. C. H. Mitchell from the University of Reading, Department of Chemistry for helpful discussions on X-ray crystallography and modelling. This work has been sup- ported by NERC. M.G.T. is grateful to the Leverhulme Trust for funding. References 1 K. Simkiss, M. G.Taylor and G. N. Greaves, J. Inorg. Biochem., 1990,39, 17. 2 G. R.Sauer and R. E. Wuthier, Bone Mineral, 1992,17, 284. 3 G. N. Greaves, K. Simkiss, M. Taylor and N. Binsted, Biochem. J., 1984, 221, 855. 4 M. G. Taylor, K. Simkiss, G. N. Greaves and J. Harries, Proc. R. SOC.London, B, 1988,234,463. 5 M. G. Taylor, G. N. Greaves and K. Simkiss, Eur. J. Biochem., 1990,192,783. 6 G. N. Greaves, S. J. Gurman, C. R. A. Catlow, A. V. Chadwick, S. Houde-Walter, C. M. B. Henderson and B. R. Dobson, Philos. Mag. A, 1991,64, 1059. 7 C. Huang and A. N. Cormack, J. Chem. Phys., 1991,95,3634. 8 B. Vessal, G. N. Greaves, P. T. Marten, A. V. Chadwick, R. Mole and S. Houde-Walter, Nature (London), 1992,356, 504. 9 A. J. de Man, B. W. H. van Beest, M. Leslie and R. A. van Santen, J. Phys. Chem., 1990,94,2524. 10 M. T. Dove, B.Winkler, M. Leslie, M. J. Harris and E. K. H. Salje, Am. Mineral., 1992,77, 244. 11 M. G. Taylor, K. Simkiss, M. G. B. Drew, P. C. H. Mitchell and M. Leslie, Mol. Simul., 1992, 9, 129. 12 G. M. Clark and R. Morley, Chem. SOC.Rev., 1976,5,269. 13 K. Lukaszewicz and R. Smajkiewicz, Rocz. Chem., 1961,35,741. 14 C. Calvo, Can. J. Chem., 1965,43, 1147. 15 T. Stefanidis and A. G. Nord, Acta Crystallogr. Sect. C, 1984, C40, 1995. 16 B. E. Robertson and C. Calvo, J. Solid State Chem., 1970, 1, 120. 17 C. Calvo, Acta Crystallogr., 1967,23, 289. 18 C. Calvo, Inorg. Chem., 1968, 7, 1345. 19 N. C. Webb, Acta Crysrallogr., 1966,21,942 20 A. B. Lidiard, J. Chem. SOC.,Faraday Trans. 2, 1989,85,341. 21 M. Leslie, Physica B (Amsterdam), 1985, 131B, 145. 22 N. F. Mott and M. J. Littleton, Trans. Faraday SOC., 1938, 34, 485. 23 P. R. Adby and M. A. H. Dampster, in Introduction to Opti-mization Methods, Chapman and Hall, London, 1974. 24 A. Hezel and S. D. Ross, Spectrochim. Acta, Part A, 1967, 23, 1583. 25 J. H. Harding and A. H. Harker, Harwell Report, AERE-R10425, 1982. 26 R. G. Gordon and Y. S. Kim, J. Chem. Phys., 1972,56,3122. 27 P.T. Wedepohl, Proc. Phys. SOC.,1967,92, 79. 28 E. Clementi and C. Roetti, At. Nucl. Data Tables, 1974, 14. 29. F. Herman and S. Skillman, in Atomic Structure Calculations, Prentice-Hall, Englewood Cliffs, 1963. 30 J. C. Slater and J. G. Kirkwood, Phys. Rev., 1931,37,682. 31 N. C. Pyper, Philos. Trans. R. SOC. London, Ser. A, 1986, 320, 107. 32 D. W. J. Cruickshank, J. Chem. SOC., 1961, 5486. 33 B. G. Dick Jr. and A. W. Overhauser, Phys. Reu. 1958,112,90. Paper 3/04862D; Received 1lth August, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000641

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(4), 649-652 Adsorption of Binary Mixtures of Heptane and Alkanols by Activated Carbon Amelia M. Gonqalves da Silva," Virgilio A. M. Soares and Jorge C. G. Calado Centro de Quimica Estrutural, Complexo I, lnstituto Superior Tecnico , 1096 Lisbon, Portugal The adsorption isotherm of heptan+hexanol on activated carbon at 298 K is presented and compared with our earlier work on heptane-ethanol and heptane-butanol mixtures under similar conditions. The data show that the preferred adsorption shifts from the alkane to the alkanol as the length of the alkanol chain, relative to that of the alkane chain, increases. The surface mole fraction, interfacial tension differences and surface activity coeffi- cients have been estimated based on an approximate theoretical model (surface phase model).These results show that the minimum thickness of the layers required to achieve thermodynamic consistency between the surface phase model and the experimental data increases as the length of the alkanol chain decreases. All the systems exhibit smaller deviations from ideality in the adsorbed phases than in the bulk. This work is part of a larger study of the adsorption of binary mixtures of alkanes and alkanols on activated carbon with the purpose of examining the applicability of thermody- namic methods of analysis, and of testing the adsorption models and their interpretation in terms of molecular behav- iour. Previously, we reported the adsorption data for the system (heptane-ethanol)/activated carbon and their interpretation by simple models.' This study was complemented with infor- mation obtained from the adsorption on the same adsorbent of the pure components from the gas phase.In another paper it was shown that the three binary mix- tures involving butanol, heptane and dodecane in contact with activated carbon behave, within the experimental preci- sion, in a mutually consistent fashion, from a thermodynamic viewpoint.2 The surface phase model was tested on all the systems and it was found that only the nearly ideal system (dodecane-heptane) shows a good agreement between theory and experiment; the remaining systems, all involving alka- nols, display strongly non-ideal behaviour and are not satis- factorily described by that simple model.This paper reports measurements on the system heptane- hexanol on activated carbon at 298 K; a comparison is made with our earlier work on mixtures of heptane-alkanol (ethanol and butanol). Experimental Adsorption isotherms were determined by the conventional immersion method, using a differential refractometer (R 401, from Waters Associates) to analyse the equilibrium composi- tion of the bulk solutions. The procedure used has already been described.2 Materials Commercial activated carbon, 'pro analysis' grade from Merck, was used without purification. Hexanol ( 2990/,) and heptane ( 299.5%) were puris. grade from Fluka, butanol ( 299.5%) and ethanol ( 2990/,) were pro analysis grade from Merck and all were used without further purification.Their purity was further checked by adding suit- able quantities of adsorbent to each liquid (in the adsorption cells) and verifying that no significant change in the refractive index of the liquids occurred after contact with the adsorbent. Physical characterization of the activated carbon was made by adsorption of nitrogen at 77 K and pure heptane and pure ethanol at 298 K, as reported previously.' The analysis of these adsorption isotherms by the BET method and the a, plot of the nitrogen isotherm revealed the existence of narrow pores of different width, including ultramicropores in which the accessibility of ethanol and heptane is restricted. The spe- cific surface areas, us (BET) obtained from nitrogen, heptane and ethanol adsorption were 860, 646 and 697 m2 g-', respectively. Results The specific surface excess of the preferentially adsorbed com- ponent, 2, can be expressed in the following form:3 n;(")fm= noAxifm (1) where no is the total amount of components 2 and 1 in the system and Ax; = xi -xi is the variation of the bulk mole fraction when the solution, of initial concentration, xi, is equilibrated with a mass m of the solid.The specific surface excess isotherm as a function of xi for the system heptane(1)-hexanol(2) at 298 K is shown in Fig. 1. The isotherm data are given in Table 1. The experi- mental data follow a type IV isotherm in the Schay and Nagy cla~sification.~The hexanol(2) is the component preferentially adsorbed for xi values up to 0.6 and heptane(1) is preferentially adsorbed at higher values of xt .The measure- ment of the adsorption at different times of contact between solution and adsorbent for similar systems showed that the adsorption equilibrium was reached in 24 h.' Moreover, the -0.2L 1 I 1 I 0 0.2 0.4 0.6 0.8 1.0 x: Fig. 1 Specific surface excess isotherm of [heptane(l)-hexanol(2)]/ activated carbon at 298 K. Symbols represent experimental points; (-) calculated by eqn. (2) and eqn. (3); (---) calculated by eqn. (2) using K = 1.59. 650 Table 1 Values of (n;(")/rn)for heptane(1)-hexanol(2) on activated carbon at 298 K 0.0180 0.449 0.5027 0.125 0.0287 0.533 0.5641 0.045 0.0325 0.510 0.6532 -0.032 0.0565 0.584 0.7639 -0.088 0.0914 0.59 1 0.8041 -0.106 0.1463 0.539 0.8468 -0.104 0.2106 0.500 0.9026 -0.097 0.298 1 0.380 0.9336 -0.068 0.4001 0.244 0.9710 -0.030 0.4550 0.190 points in Fig.1 for the various concentrations were obtained with different contact times, ranging from 24 to 72 h; the very low scatter of the experimental points is a further confirma- tion of the true equilibrium nature of the data. Discussion Comparison with Other Alkane-Alkanol Systems One particularly interesting feature of the alkane-alkanol mixtures adsorbed by activated carbon is that the preferential adsorption shifts from the alkane to the alkanol as the length of the alkanol chain, relative to that of the alkane chain, increases.This is illustrated in Fig. 2 for three mixtures with heptane : heptane is preferentially adsorbed from mixtures with ethanol and butanol in the composition range up to a mole fraction of heptane of 0.6,while hexanol is preferentially adsorbed from mixtures with heptane up to a mole fraction of hexanol of 0.6. These results follow a similar trend to that obtained on graphon by Everett.' At the same temperature (298 K), heptane is preferentially adsorbed from mixtures with ethanol and butanol over the whole composition range, while hexanol and octanol are preferentially adsorbed from the mixtures with heptane. Everett suggested that the preferential adsorption depends on the ratio of the number of C and 0 atoms in the two molecules which can interact simulta-neously with the surface.These results suggest that the long chain of the alkanol lies roughly parallel to the surface. This orientation is mostly determined by the dispersion forces between the long hydrocarbon chains and the surface and is probably enhanced by the hydrogen bonding between the OH groups. 2.0 I I I I (a1 ---C v -0.5 -J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 The systems presented in Fig. 2 also show that the adsorp- tion (selectivity) increases with the difference in the chain lengths of the two components, a fact which is consistent with the view that the adsorption arises mainly from dispersion forces between alkyl groups and the carbon surface. The S shape of the adsorption isotherms exhibited by heptane-alkanol mixtures is clearly related to the surface polarity of the activated carbon, since the adsorption iso- therms of the same mixtures on graphon (with a non-polar surface) are U-shaped.Surface Phase Model The surface phase model6 provides a relation between the specific surface excess and the composition of the bulk liquid where K is the equilibrium constant of the phase-exchange reaction between a binary solution of components 1 and 2 in contact with a solid adsorbent, ns is the number of moles in the adsorbed phase, and yi,7; are the activity coefficients of components 1 and 2 in the bulk solution. Eqn. (2) assumes an ideal adsorbed phase of equal size molecules. K is related to the interfacial tension difference, as follows K = exp[-(at -af)u,/RT] (3) where a:, a; are the interfacial tension between the pure liquids, € and 2, and the activated carbon, a, is the molar area occupied by component 1 in the interface, a, = uJns, with usbeing the specific surface area of the solid.In order to apply eqn. (2) to the experimental results we need to evaluate (a; -a:)uJRT by integration of the Gibbs equation from x;=otox;= 1: Since there are no reliable data available for the system (heptane-hexanol)/activated carbon the bulk activity coeffi- cients were estimated using the UNIFAC method.' The curve of (n;(")/m)/(xix:7;) vs. xi 7; shown in Fig. 3 was determined using both the experimental points and inter- polated ones. Integration procedures, making use of the Newton Cotes formulae have been described before.' The amount of substance in the adsorbed phase, ns = 1.81 k0.06 mmol g-', was determined through the best 4 I 1 I I -3-I 8El, -2-8 -f .8 -"8 8 .-ii--------W 8 -1 -'8 ..I I I I J-2 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 fit of eqn. (2) and eqn. (3) to the experimental isotherm. The result of the integration of eqn. (4), (a; -at)uJRT = -1.06 0.02 mmol g-' was then used in eqn. (3). The agreement between eqn. (2) (Fig. 1, solid line) and the experimental data (symbols) is quite good. The K value can also be estimated from the azeotropic 2point, since at the azeotropic composition x; = xi, nu(")= 0. Assuming an ideal adsorbed phase of equally sized molecules, K becomes equal to yi/yi.The best fit of eqn. (2) with K = 1.59 is represented in Fig. 1 by the broken line. The agreement between the model and the experimental results is worse than before, and this can probably be attributed to the approximate estimation of the bulk activity coefficient data. Surface Mole Fraction Assuming the surface phase model, Everett6 has derived the following equation to calculate the surface composition of an adsorbed phase of z layers of molecules on a plane, smooth, completely covered homogeneous surface : where a: and a! are the areas occupied by the molecules in a monolayer. The adsorption results, for the system (heptane- hexanol)/activated carbon, are compatible with a monolayer because, for z = 1, x; d 1 and xi increases with xi over the whole composition range.This test was applied to the present system using molecular areas derived from vapour adsorp- tion : a:(heptane) = 0.568 nm2 molecule-' (experimental value determined by Clint' for close-packed molecules with the major axis parallel to the surface of the graphon) and ai(hexano1) = 0.414 nm2 molecule- (estimated value). The latter value was estimated from data for other alkanols, since no experimental data could be found in the literature for hexanol. Given the fact that the OH group in alkanols allows for different orientations at the surface, we have adopted for each alkanol molecule the average value, calculated from a set of experimental values on several adsorbents;' these average values were then plotted as a function of the number of the atoms of carbon in the chain; the a! for hexanol is then obtained through interpolation for n = 6.The a, value used was 646 m2g-obtained by heptane vapour adsorption. This test has been applied before to the systems heptane- ethanol' and heptane-butanol on activated carbon' and it was found then that the minimum thicknesses of the surface layer which obeys the above conditions (xi < 1 and x; increasing with xi over the whole range of composition) are z = 3 and z = 2, respectively. For all three heptane-alkanol systems, we can conclude that the minimum thickness of the surface layer (zmin)decreases when the chain length of the alkanol increases. Interfacial Tension Difference The value of (a -a:)aJRT, where D is the surface tension of a solution of mole fraction xi in contact with the solid, derives from eqn.(4) where the integration limits are now xi = 1 and xi . The combination of eqn. (4) with eqn. (2) pro-vides a relation between the interfacial tension and the com- position of the bulk liquid (a -a*)a = -ln{xiyi + xiy: exp[(o; -of)a,/nsRRT]) (6)n'RT Comparison of eqn. (4) with eqn. (6) in Fig. 4 shows that there is good agreement between the experimental results and 65 1 I I I (a) ' 1.0 I--U (Dh +N I b Y 1.0 1 r -lS *-2.0 0 0.2 0.4 0.6 0.8 1.0 x: 1.5 1.o c I 0 0.5 0 -0.5 -1 .o -1.5 I I 1 I J-2.0 I 0 0.2 0.4 0.6 0.8 1.0 x: Fig.4 [(a -at)aJRT] as a function of xi, (m) calculated from eqn. (4) and (V)from eqn. (6), for: (a) heptane(l)-hexanol(2);(b)heptane(2)-butanol(1); (c) heptane(l)+thanol(2) on activated carbon the theory, except for the system heptane-thanol. This is probably due to the fact that eqn. (6) is based on the assump- tion that the molecules have equal sizes, which can hardly be true in the case of the heptane-ethanol system. Surface Activity Coefficient The activity coefficients in the adsorbed phase, y;, can be evaluated through the following equation where xi is given by eqn. (5), with z = 1 for heptane-hexanol, T = 2 for heptane-butanol, and z = 3 for heptane-ethanol and (a -at)a$RT obtained from eqn.(4). The activity coeff- cients at the interface yt are plotted us. xi in Fig. 5. The curves of 7: us. xi are also shown in the same figure. The present data together with those obtained for the other two systems seem 652 1 I I 8 6 x 4 2 0 0.2 0.4 0.6 0.8 1.0 4 12 I I IB' I\ ~ 0 0.2 0.4 0.6 0.8 1.0 4 Fig. 5 Activity coefficients in bulk phase (---) and in adsorbed phase (-) as a function of bulk composition: A heptane(1)- hexanol(2); B heptane(2)-ethanol(l) on activated carbon. (a)y:, (b) y:,(4I4and (4Ys2. to indicate that the adsorbed phases of the mixtures of alkanes and alkanols on activated carbons exhibit smaller deviations from ideality than the bulk phases. Similar conclu- sions were reached by Kiselev and Khopina," and Nagy and J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Schay'' from results obtained with strongly non-ideal bulk phases. Our study confirms that adsorption from solution can often be discussed satisfactorily by assuming that the adsorbed phase behaves like an ideal monolayer in equi- librium with a non-ideal bulk pha~e.'~.'~This model is acceptable when the molecules of both components are of similar size, as assumed in eqn. (2), (heptane-hexanol, Fig. 1 and 4).The discrepancies between the theoretical and experi- mental results increase with the difference in the size of the component molecules. References 1 A. M. Gongalves da Silva, V. A. M. Soares, J. C. G. Calado and M. Brotas de Carvalho, J. Chem. SOC.,Faraday Trans., 1991,87, 3799. 2 A. M. Gongalves da Silva, V. A. M. Soares and J. C. G. Calado, J. Chem. SOC., Faraday Trans., 1991,87,755. 3 Manual ofSymbols and Terminology, IUPAC, Phys. Chem. Div., Appendix 11, part 1; D. H. Everett, Pure Appl. Chem., 1972, 31, 579. 4 G. Foti, L. G. Nagy and G. Schay, Acta Chim. Hung., 1974,80, 25. 5 D. H. Everett, Progr. Colloid Polym. Sci., 1978,65, 103. 6 D. H. Everett, Pure Appl. Chem., 1986,58,967. 7 A. Fredenslund and P. Rasmussen, Fluid Phase Equilib., 1985, 24,115. 8 J. H. Clint, Faraday Trans. I, 1972,2239. 9 A. L. McClellan and H. F. Harnsberger, J. Colloid Interface Sci., 1967,23, 577. 10 A. V. Kiselev and V. V. Khopina, Trans. Faraday SOC., 1969,65, 1936. 11 L. G. Nagy and G. Schay, Acta Chim. Acad. Sci. Hung., 1963,39, 365. 12 D. H. Everett, Trans. Faraday Soc., 1964, 60, 1803; 1965, 61, 2478. 13 S. Sircar and A. L. Myers, J. Phys. Chem., 1970,74,2828. Paper 3/06096I; Received 12th October, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000649

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(4), 653-657 IR Studies of Cerium Dioxide: Influence of Impurities and Defects F. Bozon-Verduraz" and A. Bensalem Laboratoire de Chimie des Materiaux Divises et Catalyse, Universite Paris 7,2,place Jussieu, 75251 Paris Cedex 05,France An infrared study of CeO, is presented which allows discrimination of the bands due to residual carbonaceous impurities from the bands arising from multiphonon absorption and electronic transitions induced by atomic defects. CO adsorption experiments show the presence of coordinatively unsaturated (cus) Ce3' or Ce4+ surface ions depending on the nature of the pretreatment. The role of plasmon-phonon coupling in the infrared examination of semiconductors is also stressed. In recent years, much effort has been devoted to the prep- aration of cerium dioxide with high surface area, to the inves- tigation of its textural and structural changes upon increasing temperature and to the study of its surface chemistry.'-5 In addition, the role of cerium dioxide in automotive exhaust catalysts has led to renewed interest in fundamental studies concerning the interaction of oxygen, hydrogen, carbon mon- oxide and carbon dioxide with this Various experi- mental techniques were involved such as infrared spectroscopy, UV-VIS diffuse reflectance, X-ray photoelec- tron spectroscopy, magnetic susceptibility,* electrical conductivity" and temperature-programmed desorption or reduction.IR spectroscopy has been used to study the nature of the species formed upon adsorption of CO and CO, (carbonates) or of oxygen (superoxide and peroxide) upon ceria either heated in uacm or reduced by However, the propensity of ceria for non-stoichiometry ' and for strongly retaining carbonate impurities brings about some uncer-tainties, and the infrared spectra of virgin ceria and of ceria treated in reducing or oxidizing atmospheres need to be re- examined, taking into account the following points: (i) Removal of the carbonate impurities by outgassing at high temperature gives rise to defects such as oxygen vacancies and Ce3+ ions.(ii) Multiphonon bands may appear in the spectral range covered by carbonate species. (iii) In reducing conditions, ceria is an n-type semiconductor and its transmis- sion decreases as it is reduced because of the large absorption due to conduction electrons.(iv) As a consequence, chemi- sorption of electron acceptors leads to a significant decrease of the background absorbance, which induces some ambi- guities when difference spectra are recorded. The aim of this paper is to throw some light on the intri- cate influence of these factors, and provide a better under- standing of the interaction of ceria with metals in transition-metal catalysts. Experimental Materials Cerium dioxide CeO, was obtained from Rh6ne Poulenc. Its specific surface area and its mean particle size after various pretreatments are presented in Table 1. Apparatus IR spectra were recorded on a Perkin-Elmer 1730 Fourier-transform spectrometer with 30 scans at 4 cm-' resolution.The spectra presented are of the sample before adsorption, in the presence of the gas phase and upon decreasing the pres- sure. The sample was pressed under ca. 300 kg cm-, into a self-supporting disc weighing ca. 30 mg ern-,. The IR cell equipped with ZnSe windows was connected to a vacuum system with P < lop5Torr. Some experiments were carried out on a grease-free vacuum line to detect any sample con- tamination by the apparatus. Pretreatment Procedures The spectral features of cerium dioxide depend sharply on the temperature and on the atmosphere of the pretreatment. All samples were submitted to a calcination in flowing oxygen at a definite temperature T,, (673 < 7JK < 1073) for 2 h before outgassing for 15 h at a fixed temperature T, (573 < T,/K < 1173).Some samples were reduced in flowing hydrogen at T,(573 d T,IK f773) for 2 h (after the calcination at T,, and before outgassing at 7J. The samples not pretreated in H, will be referred to as unreduced. Results and Discussion Influence of the Pretreatment Procedure on the Spectrum of Ceria Because of its basic character, ceria strongly binds carbonate entities. Before performing any CO or CO, adsorption experiments, various bands due to these entities appear in the IR spectrum. The positions of the different bands observed after each type of pretreatment are collected in Table 2. U nreduced Samples After outgassing at 573 K for 15 h (Fig.1, Table 2), the spec- trum of the sample showed: (i) two intense bands at around 1460 cm-' (K) and 1390-1365 cm-' (L), (ii) two bands at 1067 and 1033 cm-' (M, N) of moderate intensity, (iii) two weak bands at 853 and 835 cm-' (P, Q) on the high- frequency tail of the phonon spectrum. After outgassing at 773 K (Fig. l), N disappeared and the above bands were less intense. Because of their thermal sta- bility, these bands must be ascribed to bulk polydentate Table 1 Textural properties of CeO, specific surface mean particle size pretreatment area/m2 g -lA 573 K, vacuum 110 100-120 173 K, hydrogen 31 1173 K, vacuum 5 1173 K, air 5 260-280 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 2 Influence of the pretreatment procedure on the spectrum of ceria sample pretreatment temperature/K band position/cm -oxidation (TJK) reduction by H, (TJK) outgassing (T,/K) J K L M N PQ 673 -573 -1460 1390-1365 1067 1033 853 835 673 -773 -1472-1454 1390-1365 1067 857 835 1073 -1073 2127 --1063 -1073 -1173 2130 -1063 -673 573 573 21 19 1466 1370 1067 -855 841 673 673 573 21 19 1478 1368 1067 -859 835 673 773 573 2127 1466 1372 1067 -859 840 species rather than to monodentate carbonate." Although remained unchanged.It follows that the major contribution many bulk polydentate structures may be envisaged, the most to M does not come from carbonate species but from ceria stable should contain three cerium-oxygen bonds, in oppo- itself. As M was not modified upon contact of oxygen (160 sition with the bridged, bidentate, monodentate or carbox- Torr) with the sample either at 293 K or at 873 K (see Fig.2, ylate species (Scheme 1).l2 later) it should not be associated with oxygen vacancies or When the calcination and outgassing temperatures were as Ce3+ ions. It is therefore proposed that this band arises from high as 1073 K, K, L, P and Q disappeared, while M multiphonon processes. This view is supported by the results obtained on a ceria monocrystal by Mochiz~ki'~ who observed a band near 1030 cm-' on the high-frequency tail of the phonon spectrum (585, 465 and 280 cm-I). This attribution does not rule out a contribution of carbonate species (see below Table 3) when absorption bands between 1600 and 1300 cm-' show the presence of these entities, i.e.when the outgassing temperature is <873 K. In addition, a very weak band, J, appeared at 2127 cm-' (Fig. l), the intensity of this band growing when T, increased to 1173 K (Fig. 2). This peculiarity shows that J cannot be ascribed to occluded CO but is associated with a defect created by the drastic outgassing temperature. This is con-firmed by the decrease of J upon O2adsorption at 293 K and its disappearance after further heating in oxygen at 873 K for 15 h (Fig. 2). Experiments performed by diffuse reflectance spectroscopy (UV-VIS) also support this view ;outgassing at 1073 K leads to the formation of Ce3+ ions which are annihi- 2000 1350 700 lated by subsequent heating in oxygen at 873 K.9914wavenum ber/cm -' Fig.1 Unreduced samples. After oxidation at T,, for 2 h and out- gassing at T, for 15 h. (a) T,, =673 K, T, =573 K; (b)T,, =673 K, T,=773 K; (c) T,, =1073 K, T,=1073 K. Q) IJ bulk polydentate bridged bidentate -~~ 1--2000 1350 700 I wavenumber/cm-' Ce" Fig. 2 Unreduced samples. Successive treatments: (a)After oxida- monodentate carboxylate tion at 1073 K for 15 h and outgassing at 1173 K for 15 h; (b)after Scheme 1 Structure of some carbonate and carboxylate species contact with 130 Torr 0, at 293 K for 15 min; (c) at 873 K for 1.5 h Table 3 Attribution of bands of adsorbed CO and carbonaceous species (1 800-700 cm -I) ____~ ~ species band position/cm -~~~~ 'bulk 'carbonate 1480-1420 1400-1350 1060 880 monodentate carbonate 1480-1460 11380-1360 1070 850 bidendate carbonate 1610-1 580 1310-1250 1030 860-830 bridged carbonate 1750-1700 1180 carboxylate 1580-1560 1400-1350 J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Samples reduced by Hydrogen In addition to the bands, K-Q, ascribed to carbonate species, the sample reduced at 573 K exhibited a very weak band, J, near 2118-2120 cm-' (Fig. 3). The intensity of J increased with T, and a shift to 2127 cm-' was noticed for T,= 773 K. This is the exact position of a band ascribed by Laachir et aL9 to occluded CO arising from the reduction by H, of car- bonate entities initially observed on a sample presenting a high surface area. According to these authors, this assignment was supported by the detection of CO in TPR experiments near 900 K. However, on a sample with a low surface area, these authors noticed that the same band grew with reduction temperature in the 873-1073 K range and is destroyed in 0, at room temperature.As long as bulk carbonates are present in the sample (absorption bands in the 1600-1300 cm-' range), the pres- ence of occluded CO cannot be discarded. However, if an occluded species could be expected to resist outgassing up to T,= 573 K, it should not be so sensitive to oxygen at room temperature. Hence we believe that the J band observed before CO adsorption in the 21 18-2127 cm-' range arises from an electronic transition from donor levels located near the conduction band such as Ce3+ ions or oxygen vacancies.Surface Ce3+ ions were detected by XPS9,'5 while the pres- ence of oxygen vacancies is deduced from electrical conduc- tivity measurements.'0*16 It is also relevant to note that a band near 2115 cm-' was already assigned to defects in cerium dioxide.6 Finally, the weakening of M when T, increases (Fig. 3) is ascribed to phonon-plasmon coupling' (see below). This coupling may also partly explain the marked intensity decrease of K and L. Oxygen Adsorption According to Li et a1.,6.'8 oxygen adsorption at 298 K on ceria outgassed at lo00 K gives rise to superoxide 0; species well characterized by v(0-0) = 1126 cm-' (checked by iso- topic experiments), but two bands, not discussed by the authors, appeared simultaneously at 1363 and 1342 cm-'.On a sample reduced by H, at 673 K and outgassed at lo00 K, a peroxide entity was detected [v(O-0) = 883 cm-'1 together with the superoxide. Aside from the bands ascribed to these entities, the spectra also showed two bands in the 1550-1200 cm-' range, not discussed by the authors, and two 'reverse bands' at 939 and 21 15 cm-' assigned to CeO, species (x < 2) present on the reference spectrum. The inten- sity of these four bands was shown to increase with the adsorption temperature between 200 and 473 K. While the 0; and 0;-entities are well characterized, the other bands 4 L KA a m e s % I I 2000 1350 700 ' wavenurnber/cm -Fig. 3 Samples reduced by H,. After oxidation at 673 K for 2 h and reduction by hydrogen for 2 h at (a) 573 K; (b)673 K; (c) 773 K.shown by Li et al. call for some comments. The peaks appearing between 1550 and 1200 cm-' upon oxygen adsorption may be ascribed to carbonates and carboxylates formed by reaction of oxygen with residual carbonaceous entities. To obtain ceria samples free of contaminants requires a pretreatment in oxygen at T,, > lo00 K, in order to burn out surface hydrocarbon impurities and occluded CO, before out- gassing at T,> lo00 K so as to remove all carbonates. This conclusion arises from experiments performed on two samples pretreated at T,, = 673 and 1073 K, respectively, and both outgassed at T,= 1073 K (Fig. 4). On the former sample, 0, adsorption gave rise to bands near 1570 and 1305 cm-',which did not appear with the latter sample.Concerning the so-called reverse bands at 2115 and 939 cm-', reported by Li et it is relevant to note that plot- ting only difference spectra may lead to difficulties, especially when oxygen adsorption is involved. In n-type semicon-ducting oxides, indeed, oxygen adsorption not only removes bands associated with defects previously created in a reducing atmosphere, but also gives rise to a marked diminu- tion of the background absorbance, due to the decrease of the concentration of conduction electrons (the oxygen species acting as electron traps). It must then be emphasized that absorption by conduction electrons is a function of vP2,l9 which implies that the decrease of the background absorb- ance upon oxygen adsorption is much more important at low wavenumbers. It follows that determining the position of absorbance maxima from difference spectra is not safe in this case and that examination of direct spectra has to be pre- ferred, at least in a preliminary way.A further complication arises from the particle size (6) dependence of the light intensity Id diffused by the sample which is expressed by: I, 2 d3v4.,0 As a consequence, Id decreases with v and the sample transmittance increases when v diminishes. Hence, it is important to check that the particle size does not vary during the course of adsorption-desorption experiments, which occurs when high pretreat- ment temperatures have been chosen.CO Adsorption Carbonate and Carboxylate Species As mentioned above, carbonate and carboxylate species are easily formed and various structures may coexist, even though their thermal stabilities are different, as discussed by 4 940 ........_.... 2000 1225 456 wavenumber/cm -' Fig. 4 Samples not reduced by H,. Influence of the calcination temperature on the IR spectrum of ceria upon oxygen adsorption (100 Torr) at 293 K. (a) T,, = 673 K, T,= 1073 K; (b)T,, = 1073 K, T,= 1073 K. Busca and Lorenzelli" in a well documented review on oxides. In fact, the nature of adsorption sites is expected to change with the atmosphere and the temperature of pretreatment. The frequency of the vibrational modes should indeed be sen- sitive to the degree of coordinative unsaturation of Ce"+ ions and to the oxidation state of cerium (+3 or +4) because of the role of the polarizing power of the metal ion.Hence the detailed assignment of all bands can only be tentative. The nature of the carbonate and carboxylate entities formed on ceria was discussed previously by several auth~rs~-~and our objective is essentially to discriminate the contribution of these species from the bands arising from ceria itself. Unreduced Samples As shown above, all the samples pretreated at T,< 773 K contain bulk polydentate carbonate species responsible for bands between 1800 and 700 cm-' (Table 2 and 3). Adsorp-tion of CO (100 Torr) at room temperature [Fig. 5 (b)] gave rise to: (i) a peak at 2165 cm- ' (H) with a shoulder (H') near 2150 cm-' assigned to linear CO adsorbed on coordinatively unsaturated Ce4+ ions, in agreement with Li et ~1.;~(ii) a set of new bands at 1607, 1546 and 1180 cm- ' ascribed mainly to carobxylate and carbonate entities (Table 3).While H and H' disappeared upon decreasing the CO pressure below 10 Torr, the other bands were stable in uucuo up to 473 K and were removed only by outgassing at 573 K for 2 h, the poly- dentate being unaffected. The behaviour of the samples pretreated at T,= 1073 K was completely different [Fig. 5(c)]. Before CO adsorption, these solids showed only the very weak band, J, at 2127 cm-',ascribed to an electronic transition (see above) and the peak, M, assigned to a multiphonon process.The poor inten- sity of J may be partly accounted for by the low surface area of aria after pretreatment at 1073 K (Table I). Upon adsorp- tion of 100 Torr CO at room temperature peak J was slightly enhanced, but no new band was detected in the 1800-700 cm-' range. It can then be concluded that many of the surface oxygen ions able to react with CO (to give the car- bonate entities) were removed during the pretreatment because of the drastic surface area decrease and of the reduction of surface Ce4+ ions to Ce3+. The formation of Ce3+ ions upon outgassing at high temperature was also shown by results obtained by diffuse refle~tance'~ and mag- netic susceptibility measurement~.~ The slight enhancement of J upon CO adsorption could be ascribed to the genesis of 1568 I1 II 2200 2100 1525 1250 975 wavenumber/cm -' Fig.5 Influence of the nature of the pretreatment on the spectrum of CO adsorbed under 100 Torr at 293 K. (-) Before CO adsorp-tion; (---) after CO adsorption. (a)Sample reduced by H, , T, = 773 K; (b)sample not reduced by H, , T,= 773 K; (c)sample not reduced byH2,T,=1073K. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 CO adsorbed on Ce3+ ions, in agreement with the wavenum- ber calculated by Zaki et al." Reduced Samples 2200-1800 cm-' Range. Before any CO adsorption, the solids reduced at T, < 673 K exhibited band J at 2119 cm-' (Table 2). Upon CO adsorption (100 Torr), a new band H (also observed on non-reduced samples) appeared at 2 165 cm-', still assigned to CO adsorbed on cus Ce4+ ions (spectrum not presented here).This band disappeared when the CO pressure was lowered to 10 Torr. For T, = 773 K, band J was observed at 2127 cm-' before CO adsorption [Fig. 5 (a)].After introduction of CO, H did not appear while J was significantly enhanced; its initial intensity was not recovered through short outgassing at room temperature, but it was restored after a prolonged evacuation (2 h) at room temperature. These results show first that the cus Ce4+ ions were transformed into cus Ce3+ ions through the H, pretreatment at 773 K. They also suggest that CO is more strongly adsorbed on the Ce3+ ions, generated in the course of the reduction pretreatment, than on the Ce4+ ions, which may be explained by a slight retrodonation from Ce3+ to co.1800-700 cm-' Range. While the K, L, M, P and Q bands initially present were enhanced, a new intense peak appeared between 1560 and 1580 cm-' together with weaker bands near 1030 and 893-900 cm-' [Fig. 5 (a)].These new bands were removed by outgassing at 573 K and are ascribed to bidentate carbonate and carboxylate entities (Table 3). Con- sidering the band intensities, the concentration of these species is much more important than in the case of non-reduced samples. In addition, the very weak bands at 1728 and 1180 cm-' show the presence of bridged carbonate entities. Conclusion The present work allows the discrimination of the bands due to residual carbonaceous impurities from the bands arising from multiphonon absorption and electronic transitions induced by atomic defects.The main results may be summarized as follows: (1) Complete elimination of carbonaceous impurities requires pretreatments at about 1073 K, first in oxygen then in uacuo. (2) The spectrum of thoroughly dehydroxylated ceria pre- sents a multiphonon band at 1063 cm- '. (3) Outgassing at temperatures d773 K creates coordi- natively unsaturated Ce4+ ions able to adsorb linear CO species vibrating near 2170-2 150 cm -(4) Outgassing at temperatures 21073 K or reducing by H, at T 2 573 K generates donor levels (Ce3+ ions or oxygen vacancies) which give rise to an electronic transition near 2120-2127 cm-' (0.26 eV).When the concentration of Ce3+ ions is large enough, (e.g. after reduction in H, at 773 K), CO adsorption gives rise to additional absorption in this range, which is ascribed to CO-Ce3+ species. (5) The distribution of carbonate-like entities generated by CO adsorption also appears to depend on the nature and the temperature of pretreatment. Carboxylate and bidentate entities are preferentially formed on samples prereduced by hydrogen. (6) Finally, it must be stressed that, on semiconducting materials, the intensity of absorption bands arising from adsorbed or bulk species depends sharply on the background absorbance. In the case of n-type semiconductors such as ceria, increasing the number of conduction electrons by reducing treatments leads to the predominance of plasmon- J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 657 phonon ~oupling'~ which hinders or even precludes the observation of purely vibrational bands. On the other hand, decreasing the concentration of conduction electrons through adsorption of an electron acceptor like oxygen allows the vibrational bands to reappear. 9 10 A. Laachir, V. Perrichon, A. Badri, J. Lamotte, E. Catherine, J. C. Lavalley, J. El Filleh, L. Hilaire, F. LeNormand, E. Quemere, G. N. Sauvion and 0. Touret, J. Chem. Soc., Faraday Trans., 1991,87, 1601. (a) R. N. Bluemental and E. K. Chang, J. Solid State Chem., 1988, 72, 330; (b) L. Eyring, Handbook ofthe Physics and Chem- istry of Rare Earths, ed. K. A. Gschneider Jr. and L. Eyring, The Rh6ne-Poulenc Company is gratefully acknowledged for the supply of cerium dioxide.11 12 North Holland, Amsterdam, 1982, Vol. 3, p. 337. G. Busca and V. Lorenzelli, Marer. Chem., 1982,7,89. K. Nakamoto, Infrared and Ramun Spectra of Inorganic and Coordination Compounds, Wiley, New York, 4th edn., 1986, References p. 252. (a)J. G. Fierro, S. Mendioroz and M. Olivan, J. Colloid Inter- face Sci., 1984, 100, 303; (b)J. L. G. Fierro, J. M. Rojo and J. M. Sanz, Colloids SurJ, 1985, 15, 75. J. L. G. Fierro, S. Mendioroz and A. M. Olivan, J. Colloid fnter- face Sci., 1985, 107, 60. J. L. G. Fierro and J. L. G. Soria, J. Solid State Chem., 1987,66, 154. J. M. Heintz and J. C. Bernier, J. Phys. C, 1986,47, 1. T. Yamaguchi, N. Ikeda, H. Hattori and K. Tanabe, J.Catal., 1981,67, 324. C. Li, K. Domen, K. Maruya and T. Onishi, J. Am. Chem. SOC., 1989,111,7683. (a) C. Li, Y.Sakata, T. Arai, K. Domen, K. Maruya and T. Onishi, J. Chem. SOC.,Faruday Trans. 1, 1989, 85, 929; (b)C. Li, Y. Sakata, T. Arai, K. Domen, K. Maruya and T. Onishi, J. Chem. Soc., Faruday Trans. I, 1989,85, 1451. A. Badri, S. Lamotte, J. C. Lavalley, A. Laachir, V. Perrichon, 0. 13 14 15 16 17 18 19 20 21 S. Mochizuki, Phys. Stat. Sol. B, 1982,114, 189. A. Rakai, A. Bensalem, J. C. Muller, D. Tessier and F. Bozon-Verduraz, 10th International Congress on Catalysis, Budapest, Elsevier, Amsterdam, 1992, p. 1875. T. Arai, K. I. Maruya, K. Domen and T. Onishi, J. Catal., 1993, 53, 117. J. M. Hermann, E. Ramaroson, J. F. Temp6re and M.F. Guil-leux, Appl. Catal., 1989,53, 117. F. Boccuzzi, C. Morterra, R. Scala and A. Zecchina, J. Chem. SOC., Faruday Trans., 1981,77,2059. C. Li, Q. Xin and X. Guo, Catal. Lett., 1992,12,297. J. T. Houghton and S. D. Smith, Infrared Physics, Oxford Uni- versity Press, Oxford, 1966. M. L. Hair, InJiared Spectroscopy in Surjbce Chemistry, Marcel Dekker, New York, 1967, p. 59. M. I. Zaki, B. Viekhaber and H. Knozinger, J. Phys. Chem., 1986,90,3 176 Tourret, G. N. Sauvion and E. Quemire, Eur. J. Solid State Inorg. Chem., 1991,28,445. Paper 3/03853J; Received 5th July, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000653

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(4), 659-665 Catalytic Reactions of o-Xylene and m-Xylene with Deuterium on Metal Films Robert J. Harpert and Charles Kernball* Department of Chemistry, University of Edinburgh, West Mains Road, Edinburgh, UK EH9 3JJ The exchange and deuteriation of o-xylene and rn-xylene have been followed by mass spectrometric analysis using evaporated metal films of iron, palladium, platinum or tungsten as catalysts, usually at temperatures in the range 273-350 K. Side-group exchange was rapid in all cases but the rates of exchange of ring atoms depended on the metal and were generally lower for atoms ortho to a methyl group. The formation of dimethylcyclohexanes was accompanied by further exchange, the nature of which varied with the metal.Significant amounts of trans-1,2-dimethylcyclohexanes were produced from o-xylene over palladium and iron. The mechanisms of the reactions are discussed in terms of dissociative adsorption for all types of exchange and a contribution of a ' roll-over' reaction of adsorbed cyclohexenes in the formation of dimethylcyclohexanes, particularly with palladium and iron. The experimental work described in this paper was carried out at the Queen's University of Belfast between 1962 and 1965.' Reactions of alkylbenzenes with deuterium over metal catalysts can provide interesting information about the rela- tive rates of exchange of different types of hydrogen atoms in the molecules and also about the nature of the addition process which can yield substituted cyclohexanes with a wide range of isotopic content.In the 1960s, results with both unsintered and sintered nickel films2 showed that the relative rates of exchange of the aromatic hydrogen atoms depend on steric rather than electronic factors and that hydrogen atoms ortho to alkyl side groups are less readily replaced. Studies with p-xylene on evaporated films of a number of metals3 showed considerable variation with the nature of the metal but at that stage there were doubts about whether the exchange of the aromatic hydrogen atoms occurred by a dis- sociative or an associative mechanism. For the xylenes, it is helpful to define symbols to describe the different kinds of hydrogen atoms in the molecules. We use S for the atoms in the methyl side groups and R,, R, and R, for the aromatic hydrogen atoms where the subscript indi- cates the numbers of methyl groups ortho to the hydrogen atom.p-Xylene3 proved to be a reactant which was relatively easy to study because it contains only two kinds of hydrogen atoms, i.e. S and R,, but for the same reason the amount of information that it can provide about the activity of the dif- ferent metals for the exchange of aromatic hydrogen atoms is limited. The case for looking at the behaviour of o-xylene and m-xylene as well is that both compounds contain more than one type of aromatic hydrogen atom in addition to methyl hydrogen atoms, in fact, rn-xylene contains all three types, R, to R,. Much of the work published in the last 30 years on the reactions of aromatic hydrocarbons with deuterium over metal catalysts has been limited to benzene.The reactions of benzene and deuterium have been examined on a range of metal catalysts such as evaporated Pd-Au alloy films,4 single crystal^,'.^ supported ni~kel,~ platinum/alumina8~9 and nickel/ZnO.'O Extensive studies on the labelling of aromatic compounds with deuterium or tritium have been carried out by Garnett and co-workers, often using heavy water as the source of the label and including homogeneous as well as heterogeneous catalysis."-' However, apart from this work, t Present address : Lorimont Enterprises BV, Hertog Hendriklaan 2,4817 JV Breda, The Netherlands. there is relatively little published information relating specifi- cally to the exchange reactions of alkylbenzenes with deute- rium on metal catalysts.Results for p-xylene have been reported for various supported Pt catalysts,16 and tert-butylbenzene and p-tert-butyltoluene have been examined over a series of evaporated metal films.17 The relative rates of exchange and deuteriation of toluene and deuterium have been determined for a wide range of evaporated metal films.18 In view of the limited amount of information in the literature, there seemed to be a case for re-examining the experimental data on the reactions of o-xylene and m-xylene with deuterium over evaporated metal films' to see whether the results provided a clearer indication of the mechanism of aromatic exchange and of the process of deuteriation.Experimental The apparatus and the technique for evaporating wires to make films have been described previo~sly.~.~ The essential feature of the apparatus was the connection of the reaction vessel (198 cm3) by a capillary leak to a Metropolitan-Vickers MS2 mass spectrometer so that continuous analysis of the reaction mixture could be made. The xylenes, 99.95% purity, were purchased from the National Chemical Laboratory, dried over molecular sieves and distilled in uacuo. The charge of hydrocarbon was 164 Pa which corresponded to 8.6 x lo'* molecules in the reaction vessel with a 22 : 1 ratio of deuterium :hydrocarbon. Gas chromatographic analyses were carried out on samples collected from the reaction vessel at the end of the reactions.The vessel was cooled in liquid nitrogen for 10 min and then the residual mixture of hydrogen and deuterium pumped off. The condensable compounds were then distilled to a removable trap, dissolved in several drops of isopentane and analysed using a 'Pye' argon chromatograph with a 1.22 m column of Celite (100-120 mesh B.S.S.) impregnated with 10 wt.% silicone oil and operated at 323 K with a flow rate of 60 cm3 min-'. Mass Spectrometric Analysis Analyses were made using low accelerating voltage, 17 eV, electrons in order to minimise fragmentation. Parent ions in the range of values of m/z from 106 to 116 were used to deter- mine the composition of the isotopic xylenes after correcting for natural isotopes and for fragmentation.The main frag- ment ions for which correction was made were those formed J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 by loss of one H or D atom from the xylenes; these ions were less than 14% of the parent ions. Corrections were made on the assumption of random loss of H or D atoms. Parent ions in the range rn/z 112-128 were used for the analysis of the dimethylcyclohexanes. There was no difficulty about the overlap of the mass ranges because negligible amounts of the lighter dimethylcyclohexanes were formed. The relative sensitivity of the mass spectrometer for xylene :dimethyl-cyclohexane was 4.95 for o-xylene and 5.50 for rn-xylene, respectively. Although the fragmentation of the dimethyl- cyclohexanes by loss of hydrogen or deuterium was only a few per cent, substantial quantities of fragment ions were formed by loss of CX,, CX, and CX, (X representing H or D).These ions amounted to typically 300, 50 and 1%, respec-tively, of the parent ions and gave peaks in the range rn/z 95-110. The size of these corrections did not cause major problems because in most cases rather small amounts of the heavier dimethylcyclohexanes were formed and it was still possible to obtain values for the lighter xylenes. With Pd, which catalysed exchange much more efficiently than addi- tion, the amounts of all the lighter xylenes had become negli- gible before the dimethylcyclohexanes were formed. Results The usual method of carrying out experiments involved fol- lowing the reactions for approximately 1 h at 273 K (or sometimes at a lower temperature) and then continuing for periods of ca.30 min at one or more higher temperatures. Initial or earliest-measurable distributions of exchanged xylenes obtained mainly at 273 K, are given in Table 1. Graphs of the yields of the various isotopic products with time, such as those shown in Fig. 1-5, were used to divide the hydrogen atoms in the molecules into groups A, B, C etc. according to the ease of exchange. With o-xylene over Pd at 273 K, Fig. 1, products up to D, were formed readily but further exchange was very slow. However reactions at higher temperatures showed that two further hydrogen atoms were replaced more rapidly than the final two.In contrast over Fe at 273 K, Fig. 2, products up to D, appeared readily with o-xylene. The results in Fig. 3 for W gave some indication that the number in the groups A :B :C were 2 : 6 :2. Fig. 4 and 5 show that, over both Pt and Fe, ready exchange of seven hydrogen atoms of rn-xylene occurred but that further exchange was even slower over Fe than over Pt. The clearest evidence that the 10th hydrogen atom of rn-xylene was not easy to exchange was obtained over Fe; at 325 K when 24% of D, had been formed there was no trace of Dlo. Small amounts of the D,, product were observed over Pd but only 60 h Y 0 10 20 30 40 50 t/min Fig. 1 The exchange of o-xylene on 15.9 mg Pd at 273 K at 373 K. The groupings of hydrogens according to rate of exchange are brought together in Table 2.In order to confirm that the least reactive hydrogen atom in m-xylene was in the position ortho 8o r I 60 h s v v) .-0 40 .-0 P 0 4-.-0, 20 to the two methyl 0 2 4 6 tlmin Fig. 2 The exchange of o-xylene on 12.3 mg Fe at 273 K Table 1 Initial or earliest-measurable distributions of exchanged xylenes at 273 K product (%) mean deuterium D5 D6 D, Dt3 D9 DIO content(MJ catalyst D, D2 D3 D4 o-xylene Pd 69 21 7 2 Fe” 13 11 14 16 Pt 77 11 6 2 W 83 10 3 2 rn-xy lene Pd 40 24 19 9 Feb 2 2 2 3 Pt 74 12 10 2 W‘ 27 21 13 11 (I Distribution after 1 min when Do had decreased to 80%.for reaction at 266 K when Dohad fallen to 88%. 0.5 ----1.41 17 10 0.4 4.01 1 0.4 0.1 --1.45 0.3 0.1 ---1.32 3 ----2.18 31 50 0.1 --6.06 0.2 ----1.47 7 7 4 1 -3.29 Distribution after 2 min when Do had decreased to 27%. ‘Distribution after 1 min J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 60'r \, 501 .-0 P I/ \ \c .-2 20 lo 0 5 10 15 20 tlmin Fig. 3 The exchange of o-xylene on 7.7 mg W at 273 K 0 10 20 30 t/min Fig. 4 The exchange of rn-xylene on 8.9 mg Pt at 273 K 8o r 0 4 8 12 tlrnin Fig. 5 The exchange of rn-xylene on 15.2 mg Fe at 273 K; the D, and D, species have been omitted for the sake of clarity 661 groups, some experiments were carried out with 1,3,5-tri-methylbenzene using an evaportated nickel film.Nickel was chosen for this test because it was known' to catalyse ring exchange of other alkylbenzenes without the complications of deuteriation. As expected a rapid exchange, 14 atoms (100 molecules)-' min-' (10 mg)-', took place at 273 K but was limited to the nine D atoms in the methyl groups. On raising the temperature to 423 K, 71% of the D, compound was formed but no more highly exchanged products were detected; exchange of the remaining three atoms occurred at 473 K. Determination of Rates of Exchange The methods for the determination of the rates of exchange for molecules containing different types of hydrogen atoms have been described previ~usly.~ Under favourable circum- stances it is possible to evaluate each of the terms in the equation, k, = kA + kB + kc (1) where k, is the initial rate of entry of deuterium atoms into 100 molecules of reactant, and k,, k, and k, are the corre- sponding rates for the various types of hydrogen atoms in the molecule.For example, with o-xylene over Pd which gives exchange of the side group atoms more rapidly than the first two ring atoms, the overall rate, k, is determined from the amounts of all isotopic products and the rate for the second group, k,, from the relative amounts of the D6-D8 products. The validity of the method has been tested by studies using computer-generated data'' and shown to be reliable provid- ed that the ratio of the rates of successive groups of atoms exceeds a factor of five, i.e.kdk, > 5 and k,/k, > 5. In the results reported pre~iously,'*~~'~*'~ an attempt was made to estimate the mean number of hydrogen atoms of each kind replaced initially for molecules undergoing routine reactions (i.e. MA,M, and M,) as well as the rates. Subsequent workZo showed that for fast reactions, i.e. with rates of exchange exceeding 2 or 3% min-', the methods of evaluating M were unreliable and tended to exaggerate the values. Some rates of reaction are given in Table 2. Arrhenius parameters are reported in Table 3 for the exchange of the various groupings of hydrogen atoms and they provide a convenient method of summarising the data obtained on rates of reaction; the errors were +2 kJ mol-' for E and f0.8for log A.It was not possible to measure with accuracy rates of exchange of the fastest group of hydrogen atoms at temperatures above 273 K and so on Arrhenius parameters were obtained for group A atoms except for Pt. Likewise, because of the rapid deuteriation over W no Arrhenius parameters could be derived for exchange. Deuteriation In all cases, deuteriation gave rise to a range of isotopic dimethylcyclohexanes containing a high average number of deuterium atoms; typical initial distributions are shown in Table 4 and some results in Table 2. A feature which was common to all the deuteriation results was that the dimethyl- cyclohexanes contained more deuterium than expected for a simple addition process, i.e. some further exchange took place during the formation of the saturated products.The relevant data to demonstrate this point are given in Table 5 which includes the average deuterium content of the xylenes from which the dimethylcyclohexanes were formed and the expected number of deuterium atoms which would have been gained for a simple addition process. Arrhenius parameters for deuteriation are given in Table 6. Measurement of the J. CHEM. SOC. FARADAY TRANS., 1994,VOL. 90 Table 2 Groupings of hydrogen atoms by rate of exchange and summary of rates for exchange and deuteriation at 273K initial rates no. of hydrogen /atoms (100molecules)-' atoms in groups min-' (10 mg)-' rate of deuteriation catalyst A B C D kA kB k/% min-' (10mg)-' 0-x ylene Pd 6 2 2" -23 0.01 (4x 10-3)* --55Fe 8 2 0.06 0.06 ---6 -Pt 10 0.11 W 2 6 2 -76' 27' 1.4 m-xy lene Pd 1 2d 1 85 0.01 (0.06)b Fe 2 1 -300 0.15 (0.01)b Pt 2 1 -18 0.8 2.6 W 1 --30' -18' " The value of kdk, was 10at 348K.Rates estimated by extrapolation from higher temperatures. Approximate rates. The value of kdk, was 5 at 323K. 'Rates at 266K. cis :trans ratios of the dimethylcyclohexanes were made after The basic mechanism for the exchange of aliphatic C-H the reactions had been followed at the upper temperatures, bonds in saturated molecules involves reversible dissociative see Table 6. The percentages of trans-1,2-dimethyl-adsorption, initiated by the formation of adsorbed alkyl rad- cyclohexane formed were 50% for Pd, 69% for Fe and 11% icals.Such a mechanism will also hold for side-group for the two other metals; the percentages of the trans-1,3-exchange with alkylbenzenes. However, for the exchange of dimethylcyclohexanes ranged from 8% for Fe to 15% for Pd. aryl C-H bonds, both dissociative and associative mecha- nisms have been proposed and the early literature has been reviewed.21 It has been arguedI8 that a dissociative mecha- Discussion nism might be slow because of the high bond-dissociation General Concepts energy of phenyl-H compared with benzyl-H for which values of 464 and 368 kJ mol-', respectively, have been given.22 But In order to discuss the present results adequately, it is neces- two other factors will influence the relative rates of disso- sary to review briefly a number of concepts that have been ciative processes.The first will be the strength of the metal- established for the catalytic behaviour of aromatic molecules carbon bond of the adsorbed dissociated species which is over metal catalysts. likely to be greater for metal-aryl than for metal-benzyl Table 3 Arrhenius parameters for exchange group of temperature E log[A/molecules temperature/K where catalyst of atoms /K /kJ mol-' s-l (10mg)-'] k = 1% min-' (10mg)-' o-x ylene Pd B 273-348 63 25.2 330 Fe B 273-323 51 23.6 319 Pt all atoms 273-298 40 23.5 251 m-x ylene Pd B 273-373 40 21.3 337 Pd C 323-373 28 18.6 426 Fe B 273-325 56 25.1 296 Pt B 273-298 44 23.5 275 Table 4 Isotopic composition of the dimethylcyclohexanes produced initially product (%) temperature catalyst /K D5 D6 D7 D8 D9 DIO Dll D12 D13 D14 D15 D16 1,2-dirnethylcyclohexanesfrom o-xylene Pd 348 -----7 23 26 21 11 8 4 --16Fe" 273 6 10 8 21 9 19 4 4 3 Pt 273 5 25 25 8 11 7 8 6 5 ---W 273 8 13 12 14 15 8 12 4 7 3 2 2 1,3-dimethylcyclohexanesfrom rn-xylene Pd 323 -----4 17 24 23 19 11 2 Fe 325 -------0.2 17 50 28 5 -Pt 273 -20 19 13 15 13 12 3 5 Wb 266 -11 23 14 13 12 10 7 7 3 a Distribution after 3min.Distribution after 2min. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 5 Evidence for exchange during the formation of the dimethylcyclohexanes average D contents temperature of xylenes for sample exchange of dimethylcyclohexanes catalyst /K converted addition with addition formed o-x ylene Pd 348 5.9" 5.1 1.5 12.5 Fe 273 3.3b 5.8 1.6 10.7 Pt 273 0 6.0 2.1 8.1 W 273 0 6.0 3.0 9.0 m-xy lene Pd 323 59' 5.1 1.8 12.8 Fe 325 7.e 5.0 2.2 14.2 Pt 27 3 0 6.0 2.5 8.5 W 266 0.4 6.0 2.6 9.0 Main xylenes were: a D, 13%, D, 82%, D, 5%. D, 54%, D, 13%, D, 25%.D, 15%, D, 79%, D, 6%. D, 19%, D, 62%, D, 18%. bonds and which will also depend on the nature of the metal. Anderson23 demonstrated that there was a correlation between the rate of exchange of ethane over various metal films and the metal-metal bond strength as measured by the latent heat of vaporisation of the metal. Tungsten with a high metal-metal bond strength was one of the most active cata- lysts and palladium one of the least active.The second factor is the activation energy involved in the dissociative adsorp- tion step involving the breaking of the C-H.bond. This will probably be less for aryl C-H than for alkyl C-H because of the close approach of the metal to the aromatic hydrocar- bon through coordination of the ring parallel to the surface. An associative mechanism for the exchange of benzene will involve the reversible formation of adsorbed cyclohexadienyl radicals but there is a complicating factor because of the likely orientation of the adsorbed species with the ring paral- lel to the surface. Two reaction paths will be needed to effect e~change.~.~'If the D atom is added to the lower side of the ring then the removal of H must take place from the upper side of the ring (or vice uersa, upper addition of D and lower removal of H).There is general agreement24-27 that the rate-determining step in the hydrogenation (or deuteriation) of benzene is the formation of adsorbed cyclohexadiene because subsequent steps will be rapid. However, stereochemical considerations are also important3p2' because the rings of all the interme- diates from adsorbed benzene to adsorbed cyclohexyl species are likely to be oriented more or less parallel to the surface and this leads to the concept of 'simple deuteriation'.2* This refers to the addition of deuterium to the aromatic com- pound without any accompanying exchange of the hydrogen atoms in the benzene and the product is mainly D6-cyclohexane.No exchange or redistribution of the 6 H atoms in the reactant occurs because addition of the D atoms is assumed to take place on the lower side of the ring of the various intermediates. Interconversion between adsorbed Species such as C6H6D, and C6H6D,+' may Well OCCUr but because of the orientation of the species no replacement of the H atoms takes place. Results for the C6H6-D2 reaction at low temperatures over a number of metal films conformed closely to this behaviour.28 For xylenes, the concept of 'simple deuteriation' implies the absence of any exchange of the four ring atoms and the formation of only cis-dimeth ylcyclohexanes. There are many example^,^',^^-^^ h owever, which show that other processes play a part in the conversion of aro-matics to the corresponding cyclohexanes as well as 'simple addition'.There is evidence for the formation of cyclohexenes as intermediate products and for xylenes, the production of some dimethylcyclohexenes can provide a route to the forma- tion of trans-dimethylcyclohexanes. The extent of interme- diate cyclohexene formation is dependent on the reactant, the temperature, the nature and the dispersion of the metal, and also on the nature of the support. However, it is important to emphasise that the actual desorption of a cyclohexene inter- mediate product is not essential to bring about ring exchange alongside deuteriation or the formation of trans-products. The concept of a 'roll-over' process which involves the 'turnover' of an adsorbed alkene on the surface and was defined in relation to the exchange of cycloalkanes with de~terium~~,~'provides an adequate mechanism for both further ring-atom exchange and the formation of trans-dimet hylcyclohexanes.Exchange The results in Table 1 and 2 show the characteristics of each metal for the exchange of the two xylenes but it is also useful Table 6 Arrhenius parameters for deuteriation catalyst temperature /K temperature/K where k = 1% min-' (10 mg)-' o-x ylene Pd 298-348 36 19.7 423 Fe 273-373 40 21.5 331 Pt 298-348 32 20.3 325 W 273-298 24 19.9 264 rn-xylene Pd 323-373 36 20.8 338 Fe Pt 325-348 273-298 (46)"33 22.0(22)" (3 34)" 256 a Approximate values. to look at the relative rates for the S, R,, R, and R, hydro- gen atoms, derived from those in Table 2 and given in Table 7.Palladium shows an exceptionally high value for the exchange ratio S/R,, i.e. good activity for side-group exchange and relatively poor catalysis even for unhindered ring positions. The other three metals give similar rates for S and R, hydrogen atoms, except that the R, atoms in o-xylene exchange more rapidly than the side-group atoms over tung- sten. Thus, the general pattern for ring exchange relative to side-group exchange is W > Pt = Fe % Pd. This sequence is consistent with a dissociative mechanism for aryl exchange assuming that the rate of reaction will be faster for metals having greater metal-carbon bond strengths which in turn will parallel metal-metal bond strengths.A recent review3* has drawn attention to the trend in surface reactivity of benzene on moving across the transition series-the sequence running from non-dissociative adsorption, through partial dissociation, to complete dissociation on W(100) at low coverage.39 Cinneide and Clarke4 claimed further evidence to support a dissociative mechanism for benzene exchange over palladium-gold alloy films. They argued that exchange and addition had to involve different mechanisms because deu- teriation fell to zero at median alloy compositions whereas exchange persisted to Pd-lean alloys. Iron is the only metal to give a high value for the exchange ratio R,/R, and the other three metals show ratios between 1 and 20.All metals, as would be expected for a dissociative mechanism, show very low activity for the exchange of the R, hydrogen atom located between the methyl groups of rn-xylene. The earlier work3 with p-xylene showed that with tungsten and platinum there was a tendency to complete the exchange of the first methyl group to react because similar amounts of D, and D3 compounds were found in the initial distributions of products. Table 1 shows that a substantial degree of multi- ple exchange occurs, particularly with iron and for rn-xylene over tungsten although these reactions are too fast to permit any deductions to be made about the mechanisms. However, the similar amounts of D, and D3 products over platinum with m-xylene, and to a lesser extent with o-xylene, provide further evidence for a tendency to complete the exchange of the first group to react.These facts suggest a possible role for an a,a-diadsorbed species, =CHC,H4CH3, in the overall mechanism of the exchange in addition to adsorbed benzyl radicals. Deuteriation Tables 4 and 5 provide evidence of further exchange during deuteriation with both reactants over all four metals but the results can be divided into two groups. With palladium and iron, side-group exchange was substantially complete before Table 7 Relative rates of exchange for different types of hydrogen atoms at 273 K catalyst react ant S/R 0 Pd o-xylene 2300 10" Pd Fe m-xylene 0-xylene 8500 ca. 1 5b lo00 Fe m-xylene ca.1 2000 Pt o-xylene ca. 1 ca. 1 Pt m-x ylene ca. 1 1:a. 20 W W o-xylene m-xvlene' 0.3 ca. 1 >3 ca. 1 a At 348 K. At 323 K. 'At 266 K. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 deuteriation began and the additional exchange which accompanied deuteriation involved replacement of ring hydrogen atoms: both of these metals gave significant amount of products in the D13 to D,, range. On the other hand, with platinum and tungsten an appreciable fraction of the exchange which accompanied deuteriation must have been associated with side-group exchange and only relatively small amounts of the high exchanged D13 to D,, products were formed. The cis :trans ratios of 1,2-dimethylcyclohexanesprovide further evidence of a difference of behaviour between palla- dium and iron, and the other two metals.High percentages of trans- 1,2-dimethylcyclohexane were observed over palladium and iron but only low percentages were formed over plati- num and tungsten. These results suggest that roll-over of the various substituted cyclohexenes, leading to both ring exchange and the formation of trans- products, occurs more readily with palladium and iron, than with platinum and tungsten. However, temperature may also be a factor since most of the results in Table 4 for palladium and iron were taken at higher temperatures than with platinum and tung- sten. Recent results4' on the exchange of methylcyclopentane with deuterium have confirmed that increase of temperature favours roll-over and that the chance of this reaction depends on the metal with the sequence being Pd > Pt > Rh.The cis : trans ratios for the 1,3-dimethylcyclohexanesare less useful because the cis-product is the more stable form and so there is less tendency for trans-formation to occur. Other comments can be made about some of the features of the distributions in Table 4. The maxima at D,, and D13 for the 1,2-dirnethylcyclohexanesformed over iron probably arise from complete exchange of one methyl group and either two or four ring atoms as well as the six D added. A similar interpretation would apply to the results for tungsten where there was a partial cut-off in the distribution after the D,, product. In most cases, the rates of deuteriation were a factor of 10 faster for rn-xylene than for o-xylene although the results for iron were different.The pattern of activity of the metals for deuteriation was W > Pt > Fe GZ Pd, which is similar to results for benzene,' and toluene.18 For most systems, the energies for deuteriation were in the range 30-40 kJ mol-'. These values are somewhat smaller than those reported for various supported palladium catalysts27 for which the activa- tion energies were in the range 46-62 kJ mol-'. The turnover frequencies for reactions over palladium evaluated at 413 K using data on the area of the films41 are 1.4 x s-l for o-xylene and 17 x s-' for rn-xylene. These are close to the values reported by Rahaman and Vannice for palladium powder of 4.9 x s-' and 27 x s-' for the two xylenes measured using pressures CQ.10-fold greater than those in the present work. Conclusions These investigations establish the relative activities of the four metals (Fe, Pd, Pt and W) for both exchange and deu- teriation of the xylenes without complications associated with the supports of supported-metal catalysts, The results for the three xylenes, taken together, provide comparisons between the rates of exchange of side-group hydrogen atoms and the rates for the various types of ring atoms, classified by the number of neighbouring methyl groups. The variations exhibited by the metals, for the rates of exchange of the different hydrogen atoms are consistent with a dissociative mechanism for aryl-H exchange.The conversion of the xylenes to dimethylcyclohexanes always involves more than a 'simple deuteriation'. Some J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 665 additional exchange of ring atoms accompanied by the for- mation of trans-1,2-dimethylcyclohexanesfrom o-xylene, par- ticularly over palladium and iron, is also found. The results are interpreted in terms of a roll-over mechanism involving the turning over of the various dimethylcyclohexenes on the catalyst surface; in some cases these may appear as interme- diate products but this is not essential to account for the nature of the dimethylcyclohexanes formed. 15 16 17 18 19 20 M. A. Long, J. L. Garnett and P. G. Williams, J. Chem. SOC., Perkin Trans.2, 1984, 2105. J. W. Hightower and C. Kemball, J. Catal., 1965,4, 363. R. J. Harper, S. Siegel and C. Kemball, J. Catal., 1966, 6, 72. C. Horrex, R. B. Moyes and R. C. Squire, in Proc. 4th Int. Congr. Catal., Akademiai Kiado, Budapest, 1971, vol. 1, pp. R. S. Dowie, P. J. Robertson and C. Kemball, J. Catal., 1974,35, 189. R. A. Rajadhyaksha, unpublished results. 332-340. We thank Dr. R. Brown and Dr. G. S. McDougall for helpful discussion. R.J.H. held a research studentship awarded by the 21 22 S. Siegel, Adv. Catal., 1966, 16, 123. D. F. McMillen and D. M. Golden, Annu. Rev. Phys. Chem., 1982, 33,493. Ministry of Education, N. Ireland and the mass spectrometer 23 J. R. Anderson, Rev. Pure Appl. Chem., 1957,7, 165. was purchased by a grant from DSIR, London.24 25 26 J. Horiuti and M. Polanyi, Trans. Faraday SOC.,1934,30, 1164. A. Farkas and L. Farkas, Trans. Faraday SOC., 1937,33,837. W. F. Madden and C. Kemball, J. Chem. SOC., 1961,302. References 27 M. V. Rahaman and M. A. Vannice, J. Catal., 1991, 127, 251, 1 2 3 4 5 6 7 8 9 10 11 12 13 R. J. Harper, Ph.D. Thesis, The Queen’s University of Belfast, 1965. E. Crawford and C. Kemball, Trans. Faraday SOC., 1962, 58, 2452. R. J. Harper and C. Kemball, in Proc. 3rd Int. Congr. Catal., ed. W. M. H. Sachtler, G. C. A. Schuit and P. Zwietering, North- Holland, Amsterdam, 1965, vol. 2, pp. 1145-1 156. A. 0.Cinneide and J. K. A. Clarke, J. Catal., 1972, 26,233. J. Massardier, G. Dalmai-Imelik and J. Barbier, C. R. Hebd. Seances Acad. Sci.Ser. C, 1976,283,257. M. Surman, S. R. Bare, P. Hofmann and D. A. King, Surf. Sci., 1983,126,349. R. van Hardeveld and F. Hartog, in Proc. 4th Int. Congr. Catal., Akademiai Kiado, Budapest, 1971, vol. 2, pp. 295-308. R. Maurel and J. Barbier, J. Chem. Phys. Phys.-Chim. Biol.,1976, 73, 995. A. Morales, J. Barbier and R. Maurel, Acta Cient. Venez., 1979, 30,377. B. S. Gudkov, Kinet. Katal., 1979, 20, 668. J. L. Garnett, Catal. Rev., 1971, 5, 229. J. L. Garnett and R. J. Hodges, J. Am. Chem. SOC., 1967, 89, 4546. J. L. Garnett and R. S. Kenyon, Aust. J. Chem., 1974, 27, 1023, 28 29 30 31 32 33 34 35 36 37 38 39 40 41 267. J. R. Anderson and C. Kemball, Adu. Catal., 1957,9, 51. F. Hartog and P. Zwietering, J. Catal., 1963,2,79. S. Siegel, V. Ku and W. Halpern, J. Catal., 1963,2, 348. S. Siegel, J. Outlaw Jr. and N. Garti, J. Catal., 1979,58, 370. M. Viniegra, G. Cordoba and R. Gomez, J. Mol. Catal., 1990, 58, 107. N. Martin, G. Cordoba, A. Lopez-Gaona and M. Viniegra, React. Kinet. Catal. Lett., 1991,44, 381. R. Gomez, G. Del Angel and V. Bertin, React. Kinet. Catal. Lett., 1991, 44, 517. G. Del Angel, V. Bertin, P. Bosch, R. Gomez and R. D. Gonza-lez, New. J. Chem., 1991, 15,643. R. L. Burwell Jr., Acc. Chem. Res., 1969,2, 289. H. A. Quinn, J. H. Graham, M. A. McKervey and J. J. Rooney, J. Catal., 1971, 23, 35. B. F. G. Johnson, J. Lewis, M. Gallup and M. Martinelli, Faraday Discuss, 1991,92, 241. A. K. Bhattacharya, J. Chem. SOC., Faraday Trans. I, 1980, 76, 126. R. Brown, A. S. Dolan, C. Kemball and G. S. McDougall, J. Chem. SOC.,Faraday Trans., 1992,88,2405. C. Kemball, Proc. R. SOC.London, Ser. A, 1952,214,413. 1033. 14 J. L. Garnett, M. A. Long, A. B. McLaren and K. P. Peterson, J. Chem. Soc., Chem. Commun., 1973,749. Paper 3/05611B ;Receiued 17th September, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000659

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(4), 667-674 Al-Pillared Saponites Part 1.-IR Studies Sophie Chevalier, Raymonde Franck, Helene Suquet, Jean-Franqois Lambed* and Denise Barthomeuf Laboratoire de Reactivite de Surface, URA 1106 CNRS, Universite Pierre et Marie Curie, 4, Place Jussieu, 75252 Paris Cedex 05, France Al-Pillared saponites have been prepared from natural Na-exchanged Ballarat saponite. Acidic properties of these saponites have been studied by IR spectroscopy using pyridine as a probe molecule. Al-Pillared saponites exhibit both Brsnsted and Lewis acidity. Lewis sites are evidenced in the parent Na-exchanged clay as well as in the pillared clays. Brsnsted acid sites are not evidenced in the parent Na-exchanged saponite but Al-pillared saponites exhibit Brsnsted acidity as shown by a pyridinium band at 1549 cm-'. Brsnsted acid sites have been related to OH groups at 3740-3720 cm-' and 3597-3594 cm-'.The same wavenumbers are evidenced for v(0H) stretchings in acidified saponite and in Al-pillared saponites. It is assumed that an acidic attack of Si-0-AI linkages which may provide Brsnsted sites occurs on the clay sheets during the pillaring process. Many attempts have been made to improve thermal stability and pore size of pillared clays.'-5 Thermal stability and pore size are significant properties in the use of pillared clays as cracking catalysts. However, acidity is also relevant for this purpose. Al-Pillared saponites, stable up to 750"C, with a pore size greater than 8 8, have been ~repared.~.~ This study has been undertaken to check their acidic properties.Pyridine adsorp- tion by the pillared clays and subsequent desorption at increasing temperatures was followed by IR spectroscopy as a test method. Experimental Starting Materials Starting materials were a Ballarat saponite provided by the Source Clays Repository of the Clay Minerals Society (University of Missouri, USA) and an A1 polycation solution containing the 'Keggin &-isomer' polycation, designated All ,, and possibly other A1 polymers. A Chlorhydrol commercial solution from the Reheis Chemical Company was used as one of the sources of A],, . This All, polycation was also pre- pared from an AlCI, solution treated with NaOH.' Catalyst Preparation (i) The Ballarat saponite was previously purified by a sedi- mentation process.The fraction of particles less than 2 pm in size was collected and then exchanged, at least three times, with 1 mol NaCl 1-' solution. The resulting sample is desig- nated SNa and used for preparations at a concentration of 5 g 1-'. The unit cell corresponds to the formula: (Si7.262 A10.7 38)[Mg5.920Fe~.~04Ti0.006Mn0.0021Na0.600ca0.028K0.006 02,(OH), .A slightly different formula was previously proposed by Post for the Ballarat ~aponite.~ According to the chemical analyses he reported, the clay is assumed to have 0.16 half-cell octahedral positions filled with Al. In our sample, however, no octahedral A1 was detected by 27Al NMR spec- troscopy.6 (ii) The Chlorhydrol commercial solution was diluted to 0.1 mol 1-' and matured for 2 h at 60°C.This initial solu- tion, designated ACH, has a pH of 4.8. Adding a few drops of concentrated NaOH raises the ACH pH to 6. Adding a 2 mol 1-' ammonium acetate solution3 to the initial ACH solution with an NH4 :A1 ratio of 15 :1 leads to a pH of 7. These pillaring solutions of ACH were added to an SNa suspension to give an Al, in SNa of 5 mmol g- ' where Al, is the concen- tration of aluminium in solution. Samples A, B and C were prepared with ACH of pH 4.8, 6 and 7, respectively. Sample D was prepared with the solution of AlCl, titrated with NaOH and designated AHY (pH = 5.0 and Al, in SNa of 5 mmol g- '). After addition of ACH or AHY to SNa, the mix- tures have been kept at 80°C for 2 h.Further experimental details have been published el~ewhere.~. ' (iii) The preparation process used for sample A was repeat- ed replacing the ACH solution by 1 mol 1-' HCl. The corres- ponding sample is designated ASNa and called 'acidified saponite'. IR characterization has been achieved for SNa, ASNa, and samples A and C. In a companion paper," an NMR study of SNa, A, B, C and D samples is reported. Thermal Treatments SNa, ASNa and samples A, ByC and D were oven-dried at 60°C for 12 h. We refer to the All,-contacted samples after drying at 60 "C as 'Al-intercalated' samples. Amounts of SNa, ASNa and samples A, B, C and D were then heated at 500°C or 750°C for 4 h. The heating rate was 36°C h-' up to 500 "C and 90 "C h- ' between 500 and 750 "C.The heated samples were designated SNa60, SNa500, SNa750, A60, A750, B500 etc. according to the temperature of the thermal treatment. Only samples stabilized by heating at tem-peratures of 500 "C or more are called 'pillared clays'. Thermal Analyses Simultaneous thermogravimety (TG) and differential thermal analysis (DTA) have been performed on a Netzsch Simulta- neous Thermal Analyzer STA 409 in a covered alumina cru- cible. The sensivity of a 100% weight loss has been fixed at 25 mg and the DTA sensitivity at 50 pV. X-Ray Difbaction Spectra (XRD) XRD reflexion spectra were recorded on a SIEMENS D500 diffractometer using CU-Ka radiation. For the pillared sapon- ites, the d(001) spacings have been determined from the 001 line maximum of intensity.N,Adsorption (BET) Before adsorption, samples were previously degassed for 1 h at 150°C under N, flow. N, adsorption measurements were performed on a Quantasorb Junior apparatus and surface areas were determined from the BET treatment. The upper limit of partial pressure of N, was fixed at 0.1-0.12 to use the linear range of the adsorption isotherm. While the use of the BET model is questionable for pillared clays,12 we provide the BET surface areas to allow easy comparison with the lit- erature. IR Spectroscopy Al-Pillared saponites can be ground to fine powders. These powdered samples, stored at room temperature, were sieved and pelletized as self-supported wafers.SNa60, SNa500 or ASNa500 cannot be pulverized by dry grinding. It was thus necessary for these samples, to prepare films from the corre- sponding colloidal solutions. It has already been shown that hydration is a reversible process for SNa heated to 700 OC.13 Wet grinding allows formation of a colloidal solution of SNa500 and ASNa500 from which films may be obtained. The films or wafers were placed in a vacuum cell specially adapted for adsorption-desorption experiments. They were heated under oxygen flow, overnight, at 450°C and then kept for 12 h at this temperature, under a Torr vacuum. After cooling to room temperature, their IR spectra were checked. Pyridine adsorption was carried out at room tem- perature for 1 h.The pyridine desorption was followed by heating for several 6 h steps, under vacuum, at increasing temperatures. IR spectra were recorded at each stage on a Perkin Elmer 1700 FTIR spectrophotometer, equipped with a Perkin Elmer 3600 data station. Results Table 1 lists the results of the sample characterization of XRD and BET surface area measurements. The fixed A1 content varies with the preparation procedure indicating that the pillaring aluminium species may be different in charge and/or degree of polymerisation according to the pH and to the dilution." Pyridine adsorption-desorption experiments for IR mea- surements have been performed on SNa500, A500, A750 and C500 samples. IR Spectra in the v(0H) Stretching Vibration Range IR spectra in the v(0H) stretching vibration range are shown in Fig.1 and the corresponding wavenumbers are listed in Table 2. The dehydrated SNa5OO saponite [Fig. l(a)] shows three absorption bands referred to as vl, v2 and v3. Their assign- ment has been known for a long time (see e.g. ref. 14). The Table 1 XRD basal spacings and BET surface areas thermal fixed A1 sample treatment/ "C content"/ mmol g-' d(001)/A S(BET!/ m2 g- SNa 500 - 12.4 36 A 60 2.7 18.7 360 500 17.7 287 750 17.2 148 B 60 3.9 18.8 373 500 17.8 303 750 17.2 212 C 60 4.5 19.0 205 500 17.8 256 750 16.9 118 D 60 2.1 18.5 339 500 17.8 332 750 16.9 67 a From the solution. J. CHEM. SOC.FARADAY TRANS., 1994, VOL. 90 3500 3000 wavenumber/cm-' Fig. 1 IR spectra in the v(0H) range of (a) SNa500 (23.1 mg) film, or of (b) A500 (27.0 mg), (c) C500 (27.5 mg) and (d) A750 (25.8 mg) self-supported wafers strongest band (v, at 3674 cm- ') is assigned to the hydroxy- group stretching vibration of the Mg3(OH) units. The electric field associated with the interlayer cations, namely Na', Ca2+ and K+,toward which some OH groups are directed, raises the wavenumber of v1 = 3716 cm-'. The very weak v3 band at 3622 cm-' is assigned to the stretching vibration of hydroxy groups belonging to one trioctahedral Mg,R(OH) unit, i.e. to one unit where cations such as R = Mn or Fe replace Mg in the octahedral layer. In the dehydrated Ballarat saponite, the average unit cell contains four OH groups, 0.66 interlayer cations (Na' or equivalent Ca2 + and K +) and 0.12 divalent atoms replacing Mg in the octahedral sites.Assuming that the molar absorp- tion of the different hydroxy groups is identical, that no coupling between OH vibrations occurs and that no contri- bution from one type of OH group is present at the same wavenumber as another type of OH group, the relative inten- sity of the observed v(0H) bands would be consistent with the cation content in the dehydrated sample. Indeed, if one interlayer cation Na' or equivalent (K', Ca2+) perturbs one OH group, the expected intensity ratio Z(v1)/Z(v2) is found to be 0.17 and the experimental value is also 0.17. Assuming a random distribution of R substituents in the trioctahedral groups, the relative intensity of v3 versus v2 may Table 2 IR spectra in the v(0H) stretching vibration range (before pyridine adsorption) sample vl/cm- v2/cm-' vJ/cm-' vdcm-' SNa5OO 3716 3674 3622 A500 3722 3674 3622 3594 c500 3730 3674 3622 3597 A750 3736 3663 ? ? J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 be calculated as the ratio 3x2(1-x)/x3 if (1 -x) and x are the atomic proportions of substituent and Mg, respectively. For 0.12 R2+ replacing Mg in one unit cell, 1(v3)/1(v2)is cal- culated to be 0.06 and the experimental value is found to be 0.08. However, this band is weak and the experimental ratio I(v~)/I(v~)is probably not very accurate.The same assignment as for SNa500 has been proposed for a Llano Na-vermiculite" pretreated at 550 "C which exhibits two bands at 3675 cm-' for v(0H) of non-perturbed OH of Mg, units and at 3708 cm-' for v(0H) perturbed by Na' cations. The Av shift between these bands is 33 cm-' while in SNa it reaches 42 cm-'.This is the sign of a stronger inter- action of Na+ with the OH groups in SNa. Some changes occur in the spectra of the dehydrated pil- lared clays A500 [Fig. l(b)] and C500 [Fig. l(~)]as compared with SNa500: (i) The v2 band of the structural hydroxy group bonded to three Mg atoms does not shift but is strongly enhanced and broadened. The sampling may account for the enhancement. It is well known that film deposits are oriented and that the OH stretching interaction of the IR beam with hydroxy ions depends on the angle of incidence of the beam with respect to the 0-H bond axis.16 In the SNa sample prepared by depo- sition from aqueous solution, the film plane corresponds to the a-b plane of the layers and the direction of propagation of the infrared beam is roughly parallel to the 0-H bond axes resulting in a weak absorption band.The significant broadening of the v2 band may be due to the frequency shift of many OH oscillators owing to the change in their environment. This broadening is stronger in the lower wavenumber range as observed with the pillared hectorites.' (ii) In the vicinity of the v1 band assigned to OH groups perturbed by Na+ or K+ cations in SNa, a shoulder is still present at 3722 cm-' in A500 [Fig.l(b)]. This shoulder, although partially hidden by the broadening in C500, can still be guessed to be around 3730 cm-' [Fig. l(c)]. (iii) The main changes occur around 3600 cm-': two peaks at 3622 and 3594 cm-' in A500 and at 3622 and 3597 cm-' in C500 are sharply marked instead of the very weak band at 3622 cm-' in SNa500. It may be concluded that the new peak at 3594-3597 cm-' referred to as v4 corresponds to a new hydroxy type created by the pillaring process. A weak shoulder in the region 3625-3590 cm-' is assigned by Occelli and Finseth' to H-bonded hydroxy groups in pillared hecto- rite catalysts. For the A750 sample, the dehydroxylation of structural OH is probably almost achieved.The DTA and TGA curves in Fig. 2 show that structural hydroxy groups of the octa- hedral layer are condensed to give water at around 800°C lo( 30 when the sample is heated in air, at a 10°C min-' heating rate. A750 was kept at 750"C, under air flow, for 4 h. It may be assumed that this treatment is long enough to cause the corresponding dehydroxylation. Therefore, there may be expected a decrease of the v(0H) band intensities. In addi- tion, the thermal treatment at 750°C has another obvious effect on the pillared clay and the self-supported wafer of A750 shows an important diffusion in the 4000-3000 cm-' range as evidenced by the baseline. The poor transmittance of this sample results in a very noisy absorbance spectrum [Fig.l(d)]. The peaks in the 3600 cm-' range can no longer be evidenced and the one near 3720 cm-' band is shifted to 3736 cm-'. ASNa500 acidified saponite exhibits in the v(0H) range the spectrum shown in Fig. 3(b). The 3716 cm-' v1 band is shifted to 3738 cm-' and the new peak v4, already observed in A500 and C500, is also present and centred at 3595 cm-'. Pyridine AdsorptiowDesorption Eflect on the v(0H) Stretching Vibration Range Spectral changes in this range are shown in Fig. 4-7. The spectrum of SNa500 is hardly transformed by pyridine adsorption: after desorption at 150"C, no frequency shifts are observed and a less than 10% intensity decrease is seen on the v1 and v2 bands [Fig. 4(b)].The former intensity of these bands is recovered by desorption at 250°C [Fig.qc)] showing the very weak acidic character of the corresponding hydroxy groups of the SNa. As shown by the d(001) value and by the surface area (Table l),the interlayer space is col-lapsed for SNa500 and pyridine does not enter this space. A500 and C500 pillared clays both behave in the same way with respect to pyridine adsorption-desorption : (i) As for SNa500, a slight decrease of the 3674 cm-' v2 band is observed after adsorption of pyridine and desorption 3674 Q) 3716C es: -Jn 3716 3594 II I I 0 200 400 600 800 1000- 40 1 3500 3000 temperaturePC waven u mber/cm -' Fig. 2 (a) TGA and (b) DTA curves of oven-dried sample A (in air, Fig. 3 IR spectra in the v(0H) range of dehydrated films of (a) heating rate 10"Cmin-;sample weight 68.80 mg) SNaSOO and (b) ASNaSOO J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 3622 P) C e % 2P I I 4c I 3500 30( wavenum ber/cm -' Fig. 4 IR spectra in the v(0H) range of SNa500 (a)before pyridine adsorption and after desorption at (b)150, (c)250 and (6) 350 "C at 150°C [Fig. 5(b) and qb)] and the former intensity is recovered by desorption at 250°C [Fig. 5(c) and qc)]. No frequency shift is observed for this band. (ii) The v1 shoulder near 3722 cm-in A500 and near 3730 cm-' in C500 disappears on pyridine adsorption and is seen again after desorption at 450°C [Fig. 5(e) and qe)]. After desorption at 520°C [Fig. 5(f) and qf), the sharpening of the v2 band allows to see that the v1 peak is flattened and broad and its maximum may be centred between 3738 and 3722 cm-', that is, significantly higher than the correspond- ing peak at 3716 cm-' in SNa500.It seems likely that OH groups responsible for this peak in the SNa are different to those in A500 and the C500 samples. (iii) The intensity of the 3622 cm-' vj band decreases slightly on pyridine adsorption and desorption at 150 "C [Fig. 5(b) and qb)] and v3 is seen again after desorption at 250°C [Fig. 5(c) and qc)]. This behaviour is the same as for the v2 band, but after desorption at 450"C, the v3 band dis- appears again because of dehydroxylation. It is known" that hydroxy groups bonded to Fe2+, for example, are the first to undergo decomposition by hydrogen loss with concomitant oxidation of octahedral Fe2+.(iv) The 3594 cm-' v4 band in A500 disappears completely on adsorption of pyridine and subsequent partial desorption at 150°C [Fig. 5(b)]. It is seen again after desorption at higher temperatures with its strongest intensity after desorp- tion at 350°C. However, it does not recover its former inten- sity and then decreases with increasing temperatures of desorption. It is still evidenced at 520°C [Fig. 5(f)]. The cor- responding OH group is obviously more acidic than the others. For C500, a similar evolution of the 3597 cm-' v4 band is seen but at lower temperatures of pyridine desorption: the strongest intensity of v4 is recovered at 250°C [Fig. qc)] and the band is completely flattened at 520 "C [Fig.qf)]. 4000 3500 3000 waven um ber/cm -' Fig. 5 IR spectra in the v(0H) range of A500 (a) before pyridine adsorption and after desorption at (b)150, (c)250, (6)350, (e) 450 and (f) 520 "C A750 behaves in an unexpected way with respect to pyri- dine adsorption (Fig. 7). A net broadening of the main peak at 3663 cm-' is observed from 150 to 250°C just as if H bonds were created. This broadening disappears at 350 "C, as is expected of H bonds. Egect in the 1700-1300 crr-' Range Fig. 8-11 present the IR spectra of SNa500, A500, C500 and A750 after pyridine adsorption and subsequent pyridine desorption at increasing temperatures, in the 1700-1 300 cm-' range. The numerical values are listed in Table 3.SNa500 after pyridine adsorption and subsequent desorp- tion at 150°C [Fig. 8(a)] shows only vibration bands assign- ed to pyridine coordinately bonded on Lewis acid sites at 1574 and 1442 cm-' and to H-bonded pyridine at 1594 and 1608 cm-'. No Brernsted acid sites (or at least very few) are evidenced. Water molecules have been totally evacuated at 500°C [Fig. l(a)] and the interlayer space is collapsed (Table 1). It is not surprising that pyridine is not sorbed much on SNa500 and that it is bound only by H bonding with surface oxygens and by nitrogen coordination to Lewis sites on layer edges. Partial desorption of pyridine at higher temperatures [Fig. 8(b) and (c)] causes a splitting of the 1442 cm-' band as fre- quently observed on oxides'gi20 smectites2' or zeolite samples.22 It has been s~ggested'~*~' that the presence of at least two groups of Lewis sites on the surface may account J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 67 1 a) C e s 9, 4000 3500 3000 wavenumber/crn -' Fig. 6 IR spectra in the v(0H) range of C500 (a) before pyridine adsorption and after desorption at (b) 150, (c) 250, (d) 350, (e) 450 and (f) 520°C for this splitting. These groups would differ in acid strength and behave differently with respect to the temperature. Simi- larly, it has been reported by Ward23 that the exchangeable cations in zeolites can also act as Lewis acid centres and give rise to similar bands as do low-coordinated aluminium ions. However, it has been recently shown in a fl zeolite that a surface change cannot account for the splitting of the band near 1450 cm-' 24 and that the decomposition of pyridine with increasing temperatures may be considered.Thermal Q) C e s 9, 1 4000 3500 300 wavenurnber/crn-' Fig. 7 IR spectra in the v(OH) range of A750 (a) before pyridine adsorption and after desorption at (b) 150, (c) 250, (d) 350 and (e) 450 "C analysis of pyridine-treated sepiolite and palyg~rskite~~ leads to the same conclusion. A500 after pyridine desorption at 150°C [Fig. 9(a)] shows bands which may be assigned not only to pyridine coordi- nately bonded on Lewis acid sites (at 1450, 1575 and 1622 cm-') and to H-bonded pyridine (at 1613 cm-') but also to pyridinium ion (at 1549 and 1640 cm-I).The corresponding Brmsted acidity is thus correlated to the presence of a new species introduced by the pillaring process. This acidity is no longer evidenced at 450"C [Fig. 9(d)]. Thus, it may be related Table 3 IR spectra in the 1700-1300 cm-' vibration range after pyridine desorption at different temperatures sample temperaturefc BPY LPY H4, LPY BPY BPy + LPy + HPy LPY SNa5OO 1 50 1608 1594 1574 1490 1442 350 1608 1448 1594 1575 1490 1436 A500 150 1640 1622 1613 1575 1549 1492 1456 1450 520 1622 1492 1456 c500 150 1640 1622 1613 1575 1549 1492 1456 1450 520 1633? 1622 1492 1456 A750 150 1640 1622 1610 1575 1549 1492 1448 450 1610 1492 1448 BPy = pyridine adsorbed on Brmsted acid site.LPy = pyridine adsorbed on Lewis acid site. HPy = pyridine adsorbed by H bonding. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 1450T I 1700 1500 1300 wavenumber/cm-' Fig. 8 IR spectra in the 1700-1300 cm-' range of SNa500 after pyridine adsorption and subsequent desorption at (a)150, (b)250 and (c) 350°C to the hydroxy group which gives a new band at 3594 cm-' in the v(0H) vibration range [Fig. 5(b)-(f)]. This new hydroxy group is acidic and behaves differently from the other OH groups. The band assigned to pyridine on Lewis acid sites is already split, at 150"C, into two components: 1450 and 1456 cm-'. At higher temperatures the 1456 cm-' component decreases more slowly than the 1450 cm-' one [Fig.9(b)-(41.C500 after pyridine desorption at 150 "C [Fig. lqa)] shows a spectrum very similar to that of A500 with vibration bands assigned to coordinatively bonded pyridine on Lewis sites at 1450, 1575 and 1622 cm-', to H-bonded pyridine at 1613 cm-' and to pyridine bonded on Brernsted sites at 1549 and 1640 cm-'. The 1450 cm-' band assigned to pyridine on Q, t e 8P Q, C -ESIf? I 1'; 0 1500 1300 wavenumber/cm-' Fig. 10 IR spectra in the 1700-1300 cm-' range of C500 after pyri- dine adsorption and subsequent desorption at (a) 150, (b)250, (c) 350, (6)450 and (e)520°C Lewis sites is not so sharply split but shifts to 1456 cm-' with increasing temperatures of desorption [Fig.lqb)-(e)]. The Lewis :Brernsted site ratio which may be drawn from 1450 : 1549 cm-' band intensity ratios is greater for C500 than for A500 and may be related to the fixed A1 content (Table 1) of the corresponding samples. A750 shows an enhanced 1610 cm-' band and rather weak bands due to pyridine on Brernsted and Lewis sites [Fig. ll(a)]. Even though the 1610 cm-' band cannot be assigned only to H-bonded pyridine, it is consistent with observations in the v(0H) vibration range. The peak assigned to pyridine coordinated on Lewis sites is unique and centred at 1448 cm-' as one of the components in SNa500. It does not shift with increasing temperatures of desorption. 1700 1500 1300 wavenumber/crn -' Fig. 9 IR spectra in the 1700-1300 cm-' range of A500 after pyri- dine adsorption and subsequent desorption at (a)150, (b)250, (c) 350, (d)450 and (e) 520 "C J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 As already mentioned, the splitting of the band assigned to pyridine coordinately bonded to Lewis acid sites is frequent. Assuming that each component of this band corresponds to a different acid site, it would be attractive to assign the 1448- 1450 cm-' component to sites on the clay layer and the second component to sites influenced by the interlayer: 1436 cm-' in SNa500 where Na' is the major cation and 1456 cm-' in C500 and A500 where the pillar is in the interlayer. However, decomposition of pyridine with increasing tem- perature cannot be discarded and further experiments would be necessary to support one of these explanations.Discussion Both Lewis and Brsnsted acid sites have usually been evi- denced in most studies on pillared clays. In the various systems investigated, it is generally agreed to impute the main source of the Lewis acidity to the pillar^.^^-^' Indeed, if Lewis acidity is evidenced in the parent clay, the pillaring obviously enhances this acidity at least by facilitating access to the interlayer space. Generally, Lewis acidity is directly related to the kind and to the amount of interlayer pillars. On the contrary, controversial conclusions are reported on the Brsnsted acidity which seems to depend on the parent clay and on parameters such as pretreatment temperature or .~~outgassing temperature.Plee et ~1 have assigned Brsnsted acidity of Al-pillared beidellite to OH groups formed by protons captured by Si-0-A1 linkages. Occelli and Tindwa3' reported a protonic acidity due to OH groups associated with ACH pillars and He et aL3' asserted that the Brsnsted acidity is provided mainly by structural OH groups. Some discrepancies may be noticed too on the Lewis :Brsnsted ratios and on the acid strengths reported. In the IR spectra, Brsnsted acidity has been related to an OH group belonging to Si-OH-A1; its stretching frequency is centred at 3440 cm- ' in Al-pillared montmorillonite and beidellite.26 This band is coupled with a pyridinium band at 1540 cm-'. Acidic OH groups have been evidenced in Al-pillared bentonite3' by a band near 3700 cm-' which disap- pears upon pyridine adsorption.According to our results on saponite, Brsnsted acidity is not evidenced on SNa while it is shown on the Al-pillared saponites: in SNa500, vl, v2 and v3 are almost insensitive to pyridine sorption ; in Al-pillared saponites, v1 (near 3740- 3720 cm-') and v4 (at 3594-3597 cm-') bands are sensitive to the pyridine sorption and may be considered as due to acidic groups. So, with pillared saponites, we found OH groups involved in pyridinium ion formation both at higher wavenumbers than structural OH groups such as in benton- ite and at lower wavenumbers than structural OH groups such as in beidellite. Hydroxy groups of the v4 band slowly disappear from 350°C [Fig.5(c) and qc)] while those of the v1 band are still evidenced at 520°C [Fig. 5(f) and 6(f)].The spectra of synthetic saponite samples expanded by SiO, * TiO, colloidal particles after air drying3, exhibit bands at 3738, 3668 and 3590 cm-' which are very close to vl, v2 and v4 in our samples. These observations suggest that the Brsnsted acidity of the OH groups responsible for v1 and v4 absorptions does not depend on the nature of pillars. More- over, if we consider the results for acidified saponite, ASNa500, prepared in the same way as A500 by replacing the ACH solution by 1 mol 1-' HCI, and sampled indentically to SNa500 [Fig. 3(b)], the presence of a band at 3594 cm-' (wavenumber of v4 in A500) and the shift of v1 from 3716 to 3738 cm-' may be immediately noticed on the spectrum.Thus, the two bands which may be considered as due to acidic OH groups, namely v1 and v4 in Al-pillared saponites, are also evidenced exactly at the same wavenumbers in the 673 acidified saponite. When Brsnsted acidity is found in a parent clay, it may be related to an acidic treatment, cf. the case of acid-washed hectorite' as compared with ACH- pillared hectorite. By treating Na-beidellite with a 0.05 mol 1-' HCI solution, Schutz et observed that Na' is exchanged by H,Of. A strong band at 3440 cm-' in the IR spectrum is attributed to the corresponding bridging OH (Al-OH-Si). This band is also present in the spectrum of pillared beidellite26.29.33*34 at the same wavenumber. On the contrary, the parent Mg-hectorite studied by Occelli3' was not acid washed and does not sorb significant amounts of pyridine, similar to SNa500, and the band of pyridinium ions at 1546 cm-' is not present while it is seen on pillared hecto- rite samples.From our experiments, we may conclude that Brsnsted acidity results from H+ attack on the clay sheets. As it is evidenced on the non-pillared sample, the attack of Al-0-Si linkages of the layer seems to be involved, as already pr~posed.~~,~~ NMR experiments which will be developed further are also consistent with this hypothesis.' ' Moreover, the spectrum in Fig. 3(b) shows that the 3738 cm-' band has obviously a different origin from that at 3716 cm-' assigned to the Na' perturbation of OH.Na' exchange by H + is not completely achieved as seen by a very weak residual absorption and by the weaker intensity of the 3594 cm-' band as compared with that of A500. in the saponite expanded by SiO, TiO, colloidal particle^,^' a band at 3738 cm-' is also observed; it is assigned by the authors to the OH of silanol groups on the pillars. The inten- sity of this band is obviously higher with SiO, pillars than with ACH pillars. Thus, attention may be paid to the assign- ment of this band in our saponite samples to OH from silanol groups of amorphous silica as frequently observed on leached clays or in many zeolite samples. The behaviour of A750 in the v(0H) range may be interpreted by the increasing of the silica amount in this sample, probably because of the beginning of layer disorganization with simultaneous water formation.However, in Occelli's samples, the 3738 cm-band does not decrease drastically on pyridine adsorption as seen for A500 and A750 (Fig. 5 and 7). In the saponite sample treated by supercritical drying,32 it is more intense than in the air-dried sample and again does not decrease on pyridine adsorption. Usually, silanols of amorphous silica are not sen- sitive to pyridine sorption. Moreover, amorphous silica has not been evidenced by 29Si NMR spectroscopy in our samples. The assumption that a similar perturbation on structural OH groups as that due to Na+ cations is created by H30+ cations with a higher shift towards high wavenum- bers is not consistent with the experimental temperature range.H30+ cations can no longer be observed at 520°C and are not evidenced in the near 1700 cm-' range. Accord- ing to the observed wavenumber, the corresponding OH groups are rather 'free' as compared with the structural ones and we may assume that they are surface groups or groups located in the sheet edges which generally represent 10% of the total OH amount. Conclusion Both Lewis and Brernsted acidity have been evidenced in the Al-pillared saponites. Lewis sites are present both in the pil- lared and in the parent clay. According to the temperature of pyridine desorption, it may be concluded that Al-pillared saponites exhibit a very strong Lewis acidity. Brsnsted acidity of Al-pillared saponites is also strong.It is linked with the appearance of two OH-stretching vibration bands near 3735 and 3695 cm-'. Pillars do not bear Brransted acidic OH 674 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 groups directly but they do facilitate the access of protons to clay sheets. Our results are in good agreement with the assumption of a protonic attack of Si-0-A1 linkages to give this acidity. It has to be taken into account that protons may be yielded 14 15 16 V. C. Farmer, in The Znfrared Spectra of Minerals, ed. V. C. Farmer, Mineralogical Society, London, 1974, p. 331. J. M. Serratosa and J. A. Rausell-Colom, Mineral Petrogr. Acta, Part A, 1985,29,409. J. M. Serratosa and W. F. Bradley, J. Phys. Chem., 1958, 62, 1164.either by hydroxonium ions in the interlayer space provided by Na exchange with HCl or by reactions during the pillaring process, especially proton release upon pillar calcination with oxide formation. Saponite is a very convenient material to support this hypothesis: it easily dehydrates and does not show Brarnsted acidity in the Na form. Al-0-Si linkages of the layers undergo acid attack by giving acidic OH groups which may be seen at the same wavenumber (3594-3597 cm-') in the spectra of acidified saponite and of pillared saponites. A new assignment of the band near 3740-3720 an-' to free hydroxy groups located on the edges may be proposed. 17 18 19 20 21 22 23 24 M. L. Occelli and D. H. Finseth, J. Catal., 1986,99, 316. J. M. Serratosa and J. A.Rausell-Colom, Mineral Petrogr. Acta, Part A, 1985,29,399. E. P. Parry, J. Catal., 1963,2, 371. D. Ballivet, D. Barthomeuf and D. Pichat, J. Chem. SOC., Faraday Trans. I, 1972,68, 1712. J. P. Rupert, W. T. Grandquist and T. J. Pinnavaia, in Chemistry of Clays and Clay Minerals, ed. A. C. D. Newman, Mineralogical Society, Monograph no. 6, Longman, London, 1987, p. 282. R. Beaumont, P. Pichat, D. Barthomeuf and Y. Trambouze, Catalysis, 1973, 1, 343. J. Ward, in Zeolite Chemistry and Catalysis, ed. J. A. Rabo, ACS Monograph 171, Washington DC, 1976, p. 225. C-J. Jia, P. Massiani and D. Barthomeuf, J. Chem. Soc., Faraday Trans., 1993,89, 3659. References 25 U. Shuali, S. Yariv, M. Steinberg, M. Muller-Vonmoos, G. Kahr and A. Rub, Clay Miner., 1991,26,497. 1 2 3 4 5 6 7 8 9 10 11 12 J.Shabtai, F. E. Massoth, M. Tokarz, G. M. Tsai, J. McCauley, Proc. 8th Znt. Congr. Catal., Dechema, Verlag Chemie, Wein- heim, Berlin, 1984, p. 735. Jingjie Guan, Enze Min, and Zhiqing Yu, Proc. 9th Znt. Congr. Catal., ed. M. J. Phillips and M. Ternan, The Chemical Institute of Canada, Calgary, 1988, p. 104. F. Figueras, Catal. Rev.-Sci. Eng., 1988,30,457. D. Tichit, F. Fajula, F. Figueras, C. Gueguen and J. Bousquet, in Advances in Fluid Catalytic Cracking: Role in Modern Rejin- ing, ed. M. L. Occelli, ACS Symposium Series 375, American Chemical Society, Washington DC, 1988, p. 237. M. L. Occelli, in Keynotes in Energy-related Catalysis, ed. S. Kaliaguine, Stud. Su$ Sci. Catal., Elsevier, Amsterdam, 1988, S. Chevalier, R. Franck, J-F.Lambert, D. Barthomeuf and H. Suquet, Proc. 7th Euroclay Conference, ed. M. Storr, K-H. Henning and P. Adolphi, European Clay Groups Association, Dresden, 1991, p. 207. S. Chevalier, Thesis, Paris, 1992. S. Schonherr, H. Gorz, D. Muller and W. Gessner, Z. Anorg. Allg. Chem., 1981, 476, 188. J. L. Post, Clays Clay Miner., 1984,32, 147. S. Chevalier, R. Franck, H. Suquet, C. Marcilly and D. Bartho- meuf, in Expanded Clays and Other Microporous Solids, ed. M. L. Occelli and H. Robson, Synthesis of Microporous Materials, Van Nostrand, Reinhold, New York, 1992, p. 32. J-F. Lambert, S. Chevalier, R. Franck, H. Suquet, and D. Bar- thomeuf, J. Chem. Soc., Faraday Trans., 1994,90,675. F. Bergaya, L. Gatineau and H. Van Damme, in Multifunctional Mesoporous Inorganic Solids, ed.C. A. C. Sequeira and M. J. vol. 35, p. 101. 26 27 28 29 30 31 32 33 34 35 D. Plee, A. Schutz, G. Poncelet and J. J. Fripiat, in Catalysis by Acids and Bases, ed. B. Imelik, C. Naccache, G. Goudurier, Y. Bentaarit and J. C. Vedrine, Elsevier, Amsterdam, 1985, p. 343. Liu Zhonghui and Sun Guida, in Zeolites, ed. B. Drzaj, S. Hocevar and S. Pejovnik, Stud. Surf. Sci. Catal., Elsevier, Amsterdam, 1985, vol. 24, p. 493. D. Tichit, F. Fajula, F. Figueras, J. Bousquet and C. Gueguen, in Catalysis by Acids and Bases, ed. B. Imelik, C. Naccache, G. Goudurier, Y. Bentaarit and J. C. Vedrine, Elsevier, Amsterdam, 1985, p. 351. G. Poncelet and A. Schutz, in Chemical Reactions in Organic and Inorganic Constrained Systems, ed. R. Setton, Reidel, Dordrecht, 1986, p. 165. M-Y. He, Z. Liu and E. Min, Catal. Today, 1988,2, 321. M. L. Occelli and R. M. Tindwa, Clays Clay Miner., 1983, 31, 22. M. L. Occelli, K. Takahama, M. Yokoyama and S. Hirao, in Expanded Clays and Other Microporous Solids, ed. M. L. Occelli and H. Robson, Synthesis of Microporous Materials, Van Nos- trand Reinhold, New York, 1992, p. 57. A. Schutz, D. Plee, F. Borg, P. Jacobs, G. Poncelet and J. J. Fripiat, Proc. Znt. Clay ConJ, Denuer, 1985, ed. L. G. Schutz, H. van Olphen and F. A. Mumpton, The Clay Minerals Society, Bloomington, Indiana, 1987, p. 305. A. Schutz, W. E. E. Stone, G. Poncelet and J. J. Fripiat, Clays Clay Miner., 1987,35, 251. M. L. Occelli, Ffth Znternational Symposium on the Scientific Bases for the Preparation of Heterogeneous Catalysts, ed. G. Poncelet, P. A. Jacobs, P. Grange and B. Delmon, Elsevier, Amsterdan, 1991, p. 287 Hudson, NATO AS1 Series, vol. 400, Kluwer, Dordrecht, 1993, 13 p. 19. M. Kawano and K. Tomita, Clays Clay Miner., 1991,39, 174. Paper 3/04608G; Received 2nd August, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000667

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(4), 675-682 675 Al-Pillared Saponites Part 2.7-NMR Studies Jean-Franqois Lambed,* Sophie Chevalier, Raymonde Franck, Helene Suquet and Denise Barthomeuf Laboratoire de Reactivite de Surface, URA 1106 CNRS Tour 54-55,2eme etage, Universite Pierre et Marie Curie 4, Place Jussieu, 75252 Paris Cedex 05,France A saponite from Ballarat was intercalated with Al polycations by four different procedures involving different sources of polycations, and stabilised by calcination (pillared). 29Si and 27AI solid-state MNR of both the inter- calated and the pillared samples have been measured and are discussed. 29Si NMR leads us to propose a pillaring mechanism, involving H+ attack of the clay tetrahedral sheet followed by AIO, tetrahedra inversion, similar to that previously proposed for beidellites. 27AI NMR reveals that the pillars are not fundamentally modified upon calcination at 500°C although they undergo reversible dehydration reactions; at 750 "C and above, a strong pillar reorganisation occurs prior to collapse of the global structure.In a companion paper,' we described the study of Al-pillared saponites by several techniques, including XRD, N, physi- sorption, thermoanalytical methods, and, mainly, IR spec- troscopy of both the solid matrix and adsorbed pyridine. We analyse here the results of solid-state NMR studies under- taken on the same materials. The importance of NMR for the characterisation of inter- calated and pillared clays stems from the lack of periodicity in the organisation of the interlayer, which precludes the use of diffraction methods.The literature on this topic up to 1988 has been reviewed by Fripiat;, since then, however, little fun- damental progress has been made while researchers concen- trated their efforts on the development of new pillared materials and intercalation procedures rather than on funda- mental understanding of the mechanisms involved. Ballarat saponite is a good material for NMR studies owing to its low structural iron content' which allows one to obtain well resolved NMR lines, not broadened by paramagnetic inter- action. Experimenta1 Starting Materials and Catalyst Preparation The mode of preparation of the samples used in this study is described in full detail in ref.1 and 3. The parent clay was a saponite from Ballarat, which was submitted to pillaring by A1 polycations from four different source solutions: A, a com- mercial Chlorhydrol solution, diluted to [Altotall= 0.1 mol 1-' and aged at 60 "C for 2 h (natural pH = 4.8); B, solution A adjusted to pH = 6 with concentrated NaOH immediately prior to use; C, solution A mixed with a 2 mol 1-' ammon-ium acetate solution (NH, :A1 = 15 : 1, final pH = 7.0); D, a solution obtained by addition of aqueous NaOH to AlCl, (final A1 concentration = 0.1 mol 1-', final pH = 5.0) and containing mainly the Keggin, Al13, polycation: this solution is denoted AHY. The resulting materials, after washing and drying, are termed intercalated samples and designated A, B, C and D, respectively.Portions of these samples were then heated to 500 or 750"C with an optimised temperature ramp. The resulting materials are called pillared samples and designated A500- D500 or A750-D750. In addition to the main experiments on Ballarat saponite, a sample of Otay montmorillonite (a dioctahedral clay with octahedral substitutions) was submitted to intercalation fol- t Part 1 :J. Chem. SOC., Faruduy Trans., 1994,90,667. lowing procedure D: the resulting sample is called MD, and the same sample calcined at 500 "C is MD500. "Si and "A1 Solid-state NMR 29Si MAS NMR spectra were recorded on an XC300 Bruker spectrometer operating at a B, field of 7 T (corresponding to a Larmor frequency of 59.62 MHz for 29Si) and in some cases on an MSL400 Bruker spectrometer (field 9.3 T, ,'Si Larmor frequency 79.5 MHz).We used a single pulse sequence with quadrature detection, a pulse length of 5 ps corresponding to a 42 flip angle, and a recycle time of 1 s (it was verified that longer recycle times did not modify the observed spectrum). The magic-angle sample rotation frequency was 3 to 4 kHz. Chemical shifts are referenced with respect to TMS at 0 ppm, with Q8M8 as a secondary standard. 29Si Spectra were decomposed into a sum of Lorentzian components for quan- tification. Contrary to initial expectations, a decomposition into Gaussian components gave poorer fits with the experi- mental data; the overall observed trends were, however, the same in both cases.27Al MAS NMR spectra were recorded on an MSL500 Bruker spectrometer operating at a B, field of 11.7 T, corre- sponding to a Larmor frequency of 130.29 MHz for 27Al, with a single pulse sequence. A short pulse-length of 0.6 ps was chosen so that the results could be quantitatively exploited [flip angle < .n/(21 + l)]." The recycle time was 0.5 s and the spinning frequency was 10 kHz or 5.5 kHz (some samples could not easily be spun at high speeds). Signal posi- tions are referenced with respect to a 0.1 mol 1-' solution of aluminium nitrate at 0 ppm, with yttrium aluminium garnet as a secondary standard. Several double rotation (DOR) spectra were also obtained on an experimental probehead for 27Al, with spinning fre- quencies of 1 kHz for the outer rotor and 4 kHz for the inner rotor.Results 29SiNMR The 29Si NMR spectrum of the starting saponite is shown in Fig. l(a). It consists of two peaks in the Q3 range, at -95.8 and -90.5 ppm. These can be attributed to silicon atoms in the tetrahedral layer: the peak at -95.8 ppm corresponds to Si with no aluminium neighbour (designated Q3-OAl) and the peak at -90.5 ppm to Si with one aluminium neighbour (designated Q3-lAl), in conformity with the results of Sanz and Serratosa.' The probability of having a silicon with two -9 5.7 I ~~ -8 0 -9 0 -100 6 Fig. 1 29Si Solid-state MAS NMR spectrum of the initial saponite SNa: (a) after drying at 60°C; (b)after calcining at 500 "C;(c)after calcining at 750 "C or three aluminium neighbours is too low to give rise to an observable peak.A decomposition of the observed signal into two Lorentz- ian lines gave intensity contributions of 71 & 1% for Q3-OAl and 29 1% for Q3-lAl. We may then calculate the Si/A1 ratio by using the well known formula? where n is the number of A1 neighbours and In denotes the intensity of the peak corresponding to Q3-nAl. This formula simplifies here into Si/Al = 3(Z0 + 11)/11, yielding Si/ A1 = 10.4, and a formula of (Si7,30A10.70)for the tetrahedral sheet, in good agreement with the chemical analysis which provides (Si7.26 A10.74). In Fig. l(b) and (c), it is shown that the spectrum of a saponite heated at 500°C does not undergo any noticeable modification, while in a saponite heated at 750"C both peaks are somewhat broadened, resulting in some loss of resolution but no change in the relative contributions of Q3-OAl and ~3-1~1.By comparison with the results on pillared saponites (cf: infra), we think that a small peak is present at -85 ppm, in the Q2region. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Fig. 2-4 represent the 29Si NMR spectra of saponites sub- mitted to A1 intercalation through procedures A, C and D, respectively, (a) immediately after intercalation and after heating at (b) 500 and (c)750 "C. The intercalated samples (i.e. those not submitted to a calcination step) show no per-ceivable modification in peak positions or relative intensities with respect to the spectrum of the initial Na-saponite.In contrast, for samples heated at 500°C [Fig. 2(b)], the peaks are only poorly resolved: the peak at -90 to -91 ppm now appears as a shoulder of the main peak at -95 to -96 ppm. A mathematical decomposition shows that the relative contribution to the signal of the -90 ppm peak has noticeably decreased, falling to 21% for sample D, 14% for sample A and 7% for sample C. Although the precision of these decompositions is estimated between f1.5 and f4%, we are confident that the decrease of the peak at -90 ppm is directly correlated to the amount of aluminium fixed in the interlayer space (see Table 1). This correlation is rather obvious on the spectra even without decomposition. The spectra of samples heated at 750°C show a more con- siderable loss of organisation, together with the appearance of a signal at -85 to -83 ppm, in the region of Q2 Si.This signal is especially prominent for sample D, which has the least fixed Al. 27~1NMR As expected, the starting saponite SNa shows only one 27A1 peak in the region corresponding to tetrahedral A1 [S,,, = +64.5 ppm, Fig. 5(a)],confirming the absence of aluminium 6 Fig. 2 29Si Solid-state MAS NMR spectrum of sample A (a) after A1 intercalation and drying at 60°C; (b)after calcining at 500°C; (c) after calcining at 750 "C J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 I I I -80 -90 -100 6 Fig. 3 29Si Solid-state MAS NMR spectrum of sample C: (a)-(c) as Fig. 2.0,spinning side bands (SSBs). in the octahedral sheet. Recording the spectrum under DOR conditions [Fig. 5(b)]resulted in a considerable narrowing of this line, indicating that its residual width under MAS is mainly due to second-order quadrupolar interactions. The spectra of samples B and D are shown in Fig. 6 and 7, respectively, showing in each case the intercalated sample, the pillared sample calcined at 500°C and the pillared sample calcined at 750°C. In addition, Tables 2 and 3 summarise Table 1 Correlation between the amount of fixed A1 and the rela- tive importance of the Q3-lAl 29Si NMR line integrated intensity fixed Al/ of Q3-lAl 29Si NMR line sample mmol (g clay)-' (Yo) SNa 0 29 (unpillared saponite) D 2.1 21 A 2.7 14 C 4.5 7 Likely error on integrated intensity (estimated from independent attempts at deconvolution): between f1.5% (unpillared saponite) and f4% (sample C, pillared).-9f.5 I 1 I I I 1-60 -80 -1 00 -120 6 Fig. 4 29Si Solid-state MAS NMR spectrum of sample D: (a)-(c) as Fig. 2. 0,SSBs. 6 4.5 I I I I I I I 1 I ' lb0 90 80 70 60 50 40 30 20 pp n 6 Fig. 5 27AI Solid-state NMR spectrum of the initial saponite SNa dried at 60°C: (a)MAS spectrum; (b) DOR spectrum. *, SSBs. Table 2 Evolution with the calcination temperature of the "A1 NMR lines intensity ratio R, R, for R, for pillared sample intercalated procedure sample 500 "C 750 "C A 1.37 1.25 0.73 B 2.23 2.18 0.61 C 3.03 2.58 0.73 D 1.13 1.44 0.21 J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 63.9 I 6.5n I I 1 I I 1 I I 120 80 40 0 4 6 Fig. 6 27Al Solid-state MAS NMR spectrum of sample B: (a)-(c) as Fig. 2. 0,SSBs. relevant features of the spectra for all four samples (A, B, C and D) submitted to the same treatments. All intercalated samples show two peaks: a rather broad, asymmetrical peak with d,,, = +5.7 to +7.4 ppm, i.e. in the region of six-coordinated Al, and a narrower, more symmetri- cal one at +63.9 to 65.5 ppm, in the region of four-coordinated Al. Quantification of the A1 NMR peaks should be possible since we used pulse lengths <.n/(21 + 1). The ratio of the integrated intensities of the two peaks, taking into account the main spinning side bands, was given as R,= Z6-coord AJ14-coordThe interpretation of R, is addressed in the Discussion section. Table 2 shows the R, evolution with the calcination temperature.After heating at 50O0C, the R, values are unaltered, although the peak widths vary somewhat in samples B and C. Peak positions do not change significantly except for a small upfield shift for both six-coordinated and four-coordinated A1 in sample D. After heating at 750°C,in contrast, R, falls to values < 1, the most dramatic decrease in six-coordinated A1 intensity occurring for sample D (prepared with 'pure' AlI3). In addi- I I I I I I I I 120 80 40 0 -4 6 Fig, 7 27Al Solid-state MAS NMR spectrum of sample D: (a)-@) as Fig. 2. 0,SSBs. tion, the position of the four-coordinated A1 peak undergoes a significant negative shift for the samples with the lower fixed amounts of Al: -4.5 ppm for sample A and -6.3 ppm for sample D.It would be very interesting to discriminate between tetra- hedral A1 in the clay layers and tetrahedral A1 in the pillar, as is possible with intercalated and pillared beidellite~.~In general, we did not clearly observe more than one signal in the region of tetrahedral A1 for MAS spectra. This result is not very surprising since the resonance of the central A1 in All, is expected at +62.5 ppm (from our own liquid-state NMR results and literature data3v8), too close to the substi- tuting A1 of saponite sheets (+ 64.5 ppm) to be discriminated. We tried to record some DOR spectra to remove second- order quadrupolar broadening and improve the resolution.Fig. 8 shows the DOR 27Al NMR spectra of samples A, A500, B500 and 0.For A, B500 and C500, no improve- ment in resolution is observed owing to severe overlap with DOR SSBs. In the case of A500, however, we can distinguish between a sharp component at +63.4 ppm and a broader one at +56 ppm. This observation is interesting if we keep in mind that Plee et d7observed a similar shoulder at +56.5 Table 3 Evolution with calcination temperature of the positions of the observed 27Al NMR peak maximum for four-coordinated A1 and six-coordinated A1 peak maximum (ppm) six-coordinated A1 four-coordinated A1 calcined calcined procedure intercalated 500"C 750 "C intercalated 500 "C 750 "C A B C D +5.8 +6.8 +5.7 +7.4 +6.2 +6.4 +6.0 +4.5 +5.4 +6.5 +4.1 +4.2 +65.5 +63.9 +63.9 +65.4 +65.0 +64.1 +64.1 +63.4 +60.5+63.9 +64.6 +57.1 J.CHEM.SOC. FARADAY TRANS., 1994, VOL. 90 4.1 I 63.5I 100 80 60 40 20 0 -20 -40 Fig. 8 "A1 DOR NMR spectra: (a)-(c) as Fig. 2 ppm in pillared beidellite and tentatively attributed it to A1 in inverted A10, tetrahedra. In view of this, we are led to think that the asymmetry observed in DOR at the low ppm side of the tetrahedral A1 peak for sample A500 may be attributable indeed to a second signal. However, we did not try to decom- pose the observed signal into its components, a notoriously difficult task for a spin 5/2 nucleus. Another result from DOR NMR, albeit a negative one, is that the broadening of the six-coordinated A1 signal is heter- ogeneous rather than homogeneous since suppression of the quadrupolar interaction does not result in significant line narrowing.In other words, the local environment of octa- hedral A1 is somewhat variable, in both intercalated and pil- lared saponites; indeed, the shape of six-coordinated A1 peaks is reminiscent of that found in glas~es.~ Discussion 29SiNMR The 29Si NMR spectra of the parent saponite confirm the stability of the layers up to at least 750°C in the starting material (the peak at -85 ppm may indicate a very slight disorganisation after 4 h at 750 "C). In contrast, the 29Sispectra of samples A500, B500 and C500 indicate that the pillaring step induces a degree of local reordering in the tetrahedral sheet.We shall be mostly con- cerned here with the decrease of the intensity ratio of the -90 ppm peak with respect to the -96 ppm peak upon heating at 500 "C. If we insist on attributing these peaks to Q3-lAl and Q3-OA1 as in the parent saponite, we would have to admit that some of the aluminium ions in the tetrahedral sheet are replaced by silicon ions, a conclusion difficult to substantiate. It seems easier to rationalise this observation in terms of a mechanism put forward by Plee et al.,' in which chemical (covalent) bonds between the pillars and layers are formed by the inversion of some TO, tetrahedra (T = A1 or Si) of the clay tetrahedral sheet to yield T-0-A1,~,,,,, linkages.There is good reason to believe that this inversion is initiated by the Brernsted acidity provided by the decomposition of the pillars, and that it occurs in the vicinity of Al-substituted tetrahedra, which are the 'weak point' of the clay layer. The mechanism favoured by Plee et al. consisted of opening of an Si-0-A1 linkage by H+ (the appearance of Si-OH-A1 groups upon H+ attack of the Si-0-A1 links during the calcination step was indeed observed in the IR study of our samples') followed by inversion of either the AIO, or the Si04 tetrahedron to give A14-coord-O-A16-coord or ~~~~~~~~coor~ ~ ~ ~ c o o r linkages. These mechanisms are sum- ~ marised in Fig. 9A(ii) and B(ii), respectively. Mechanism A($ is the one favoured by Plee.Fig. 9B(i) illustrates the inversion of an SiO, tetrahedron with no A1 neighbours. We have included this mechanism because it cannot be completely excluded in view of prelimi- nary data obtained by us on pillared montmorillonites. Fig. 10 indicates that while the 29Si spectrum of intercalated mont- morillonite, MD, looks like that of the starting Na-montmorillonite, in the pillared sample, MD500, a large upfield shift and broadening of the resonance is obtained, indicative of strong local modifications in the tetrahedral layer upon pillaring. A similar shift was obtained by Butruille et a!. for pillared fluorhectorite,lO*" at some variance with previous results from the same tearn.l2 We may try to discriminate between these mechanisms by assessing their effects on the 29Si NMR signal and comparing with the experimental data.First, let us notice that in each mechanism, all Si remain Q3, and their number of A1 neigh- bours does not change. However, some of the Si-0-T tetrahedral angles do change, and they have a relevance for the 29Si chemical shift.I3 If we consider a given tetrahedron and assume it is inverted in the pillaring process, it is easy to see from Fig. 9 A s --..._ Interlayer spa-Fig. 9 Illustration of possible pillaring mechanisms (see discussion in the text): A, Preferential inversion of A10, tetrahedra. (i) Effect on Q3-OAl SiO, tetrahedron: unchanged. (ii) Effect on Q3-1A1 SiO, tetrahedron. B, Preferential inversion of SiO, tetrahedra. (i) Inversion of a Q3-OAl tetrahedron.(ii) Inversion of a Q3-lAI tetrahedron. that all three T-0-T angles formed with its neighbours will increase in the process. In contrast, if we suppose that it is one of its neighbouring tetrahedra that is inverted in the process, only one out of three T-0-T angles will increase while the others remain constant (as a first approximation, neglecting relaxation effects). The most precise and complete data on the correlation between average T-0-T angle and 29Si chemical shift published in the literature for tetrahedrally substituted phyl- losilicates were obtained by Sanz and Robert' working on the Na+-saturated form of synthetic saponites (also relevant is Weiss et all4). By extrapolating the data of Sanz and Robert, inversion of a given SiO, tetrahedron with no modi- fication of the first neighbours should yield a downfield shift of ca.-15 ppm on the resonance position of its central atom. In fact, in our case, the purely geometrical modification would have to be accompanied by a change in first neigh- bours (appearance of a Si-0-A1 linkage with the inter- calated polycation) having an opposite effect on the chemical shift (generally ca. +5 to +7 ppm in zeolites and clays): we thus estimate that the net chemical shift modification would be in the range -8 to -10 ppm, i.e. it should easily be obser- vable. On the other hand, inversion of a TO, tetrahedron neigh- bouring the concerned SiO, should only yield a -5 ppm shift since only one out of its three T-0-T angles is modi- fied in first approximation.Now if mechanism B(i) was operating, some isolated Q3-OAl Si would undergo a -8 to -10 ppm chemical shift, while their neighbours (also Q3-OAl but approximately three times more numerous) would undergo a -5 ppm shift. Since no signal is observed in the foreseen region, mechanism B(i) can be discarded here: i.e. in saponite, no SiO, tetrahedra without A1 neighbours are inverted on pillaring. Mechanism B(ii), preferential inversion of SiO, tetrahedra with one A1 neighbour, Q3-lAl Si, would result in a -8 to -10 ppm downfield shift of some Q3-lAl, but also in a -5 ppm shift of twice the amount of Q3-OAl, the Si labelled with a star in Fig. 9B(ii). This does not seem to correspond with experimental results either.Mechanism A($ would only result in a -5 ppm upfield shift of some of the Q3-lAl, those neighbouring the AlO, tetrahedra that effectively form bonds with the pillars. They would be shifted to ca. -95.5 ppm, too close to the Q3-OAl to be resolved from it. All Q3-OAl Si, and those Q3-lAl Si neighbouring non-linking AlO, tetrahedra, would be unaf- fected: the net result would be an apparent decrease of the initial Q3-lAl peak at -91 ppm compensated by an increase of the peak at -96 ppm. This decrease would be more pro- nounced for samples with high amounts of fixed A1 because of an elevated number of pillar-layer links. It can be seen that mechanism A(ii) is the only one compat- ible with the observed evolution of the 29Si NMR spectra.In particular, it is interesting that there is a positive correlation between the amount of A1 fixed in the interlayer space and the decrease of the -91 ppm contribution (Table 1). Samples with high amounts of fixed A1 were obtained from Chlorhydrol and may contain polycations more condensed than However, we do not expect the bonding between these higher polymers and the tetrahedral sheet to be very different from the All3 case. Therefore, we conclude that the formation of bonds between the pillars and the layers in pillared saponites occurs through the specific inversion of AlO, tetrahedra rather than of SiO, tetrahedra, in confor- mity with the hypothesis of Plee et al.' for beidellite. This conclusion is of course limited to pillared clays containing A1 in their tetrahedral sheet.For octahedrally substituted clays such as montmorillonite, the commonly held opinion is that J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 the pillar-layer interaction remains rather weak and of non- covalent nature even after calcining at high temperatures ; however, our data indicate that heating an All ,-intercalated montmorillonite at 500 "C involves considerable modifi-cations of the tetrahedral sheet with changes in the short- range order (Fig. 10). A more precise study is underway on the effect of pillaring on octahedrally substituted clays. The 29Si signals of pillared clays calcined at 750°C give information on the stability of the layers under severe thermal treatments.For the low pillar-density D750, it can be seen that the layers structure has been partly destroyed, with the appearance of a strong signal at -83.4 ppm probably corresponding to a new phase: at the same time, the BET 'surface area' collapses. ' In contrast, in both the unpillared saponite and pillared samples with higher pillar-density, the 29Si NMR signals are hardly modified with respect to the 500°C calcined samples. There seems to exist a critical pillars concentration that causes instability of the pillar-clay layer system: the reason for this critical concentration cannot be inferred from the data presented here although it is possible that variations in layer flexibility are involved. 27~1NMR As stated earlier, interpretation of the chemical shifts of tetra- hedral A1 is hampered by the accidental coincidence of the 6,,, for A1 in the saponite sheet and in the pillar.However, the DOR spectrum of pillared sample A500 yields support to the supposition that mechanism A(ii) is operative here since it shows the signal at +56 ppm attributed by Plee to inverted AlO, tetrahedra. Some remarks can be made on the evolution of the inten- sity ratio, R,, defined in the Results section. Octahedral A1 belongs only to the intercalated polycation; tetrahedral A1 I-80 -9b -A0 -1lo 6 Fig. 10 29Si Solid-state MAS NMR spectrum of: (a)Otay montmo- rillonite; (b)intercalated Otay montmorillonite dried at 60 "C (MD); (c)pillared Otay montmorillonite (MD500) J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 has two components, one corresponding to the (constant) amount of A1 in the saponite tetrahedral sheet, while the other contribution stems from tetrahedral A1 in the poly- cation and will of course depend on the number of inter- calated polycations. We can attempt to predict the NMR peak intensity ratio from the amount of fixed A1 (from chemi- cal analysis) if we make a hypothesis on the stoichiometry of the intercalated polycation. If we are dealing exclusively with Al13, each polycation will provide 12 octahedral Als, and one tetrahedral Al. It is then easy to see that the predicted inten- sity ratio, which we will call Ran, is: Ran = 12X/(1 + X) where X is the number of intercalated polycations normalised to the number of substituting A1 in the sheet, or X = x/y = (number of pillars per unit cell)/(number of substituted A1 per unit cell); y is known to be 0.70, while x can be imme- diately obtained from chemical analysis.In fact, it is consistently observed that Ran> R,. A similar observation was made by Schoonheydt16 who raised the possibility that the assumed stoichiometry for the inter-calated polycation (A14-coord)l(A16-coord)lcould be incorrect : the polycation would then have to contain more tetrahedral A1 related to octahedral A1 than was supposed. R, is not expected to be exactly equal to the molar ratio between the two types of A1 since some intensity may be lost to the satellite transitions.The question has been addressed in detail by Massiot et a1.,17 who designed tables of correc- tion factors applicable to the quantification of NMR data on quadrupolar nuclei for various values of the MAS rotational frequency (arf)and the quadrupolar frequency of the observed nucleus (aQ).Taking as a basis the apparent line- width for six-coordinated Al, we would favour a rather low value for the quadrupolar coupling constant (QCC) and thus for oQ:in this case, the corrections to R, due to intensity losses to the satellites would not be over 10% and the cor- rected R, would still be significantly too low with respect to Ran. However, basic aluminium sulfate which contains free [All,] ions shows a clear quadrupolar lineshape for the six- coordinated Als, and allows one to obtain for the QCC a rather high value of 9.8 to 10.5 kHz." One does not expect that the environment of the six-coordinated Als should gain in symmetry upon intercalation; therefore, we suspect that the QCC of six-coordinated A1 might be much more impor- tant than a cursory examination of the NMR signal seems to indicate.The singularities of the quadrupolar lineshape would then be masked by extreme broadening (homogeneous and/or heterogeneous) and the discrepancy between R, and Ran would be much reduced. Furthermore, the values obtained from DOR spectra for R, are higher than those from MAS spectra, and they are rather close to the Ran. This leads us to believe that the NMR intensity ratios do not demonstrate conclusively that aluminium species different from All, are present in the interlayer space.Thus, even though the quantitative data on 27Al NMR of the pillars are not conclusive at the moment, they have the merit of raising the question of the precise stoichiometry of the intercalated polycation in the various preparation condi- tions. Future studies will have to focus on the precise deter- mination of NMR parameters such as the quadrupolar shift 6,, the quadrupolar frequency aQand the asymmetry parameter qQ. Potentially, the evolution of these parameters in going from the isolated pillar (in such model compounds as All, sulfate) to the intercalated pillar might reveal a lot about the layer-pillar interaction. If we compare samples A, B, C and D and plot R, as a function of the amount of A1 fixed in the intercalation step, it seems to vary linearly up to a maximum of 4 mmol A1 (g saponite)-(samples D, A and B).For the highest amount of fixed aluminium (sample C), there is a deviation from linear- 68 1 it^.^ This means that in a first approximation, the species intercalated is the same in D, A and B: the effect of the higher pH value for sample B (6.0) would then essentially be to lower the positive charge on the polycation and thus allow intercalation of a higher amount of polycations to compen- sate for the constant cation-exchange capacity (CEC). The equilibrium form for aluminium is not expected to be the All, ion at this pH value, but the kinetics of further polym- erisation is rather slow" compared to the timescale of the intercalation procedure.In contrast, for sample C, a different phenomenon occurs, which may be cation polycondensation or surface precipi- tation of amorphous alumina on the clay tactoids (in rather low amounts since no alumina phase was observed by elec- tron microscopy). When considering the spectra of pillared samples (A500- D500), no significant variation is observed either in the R, value (Table 2) or in the position of the octahedral peak (Table 3, the position of the tetrahedral peak is not very informative as stated earlier). Apparently, the aluminium pillars have very much the same structure as the initially intercalated polycation. This is puzzling since All should lose all of its H20 ligands, and even condense its OH ligands, well before 500"C,8 a fact which seems confirmed by the TGA-DTA of our intercalated sample^.^ Both phenomena should result in changes in A1 coordination, easily detectable by NMR, which are, in fact, not observed.One possibility is that the pillars' Als do indeed lose some of their ligands upon heating at high temperature, but quickly regain them on exposure to room humidity before the spectrum is recorded. However, the spectra clearly rule out the transformation of the pillars to alumina which was sometimes proposed in the first studies on pillared clays, as already pointed out by Fripiat.2 In other words, an equation of the type: 13/x[A1,0,]'3x-2y)+ + zH' + (24 -~/2)H20 (1) where z = 7 -(13/x)(3x -2y), can still probably be written for the polycation calcination process, but one cannot assume .x = 2, y = 3 as written before.Most probably, the pillar nuclearity does not change and thus x = 13. The stoichiom- etry of eqn. (1) has received surprisingly little attention so far, and we are currently conducting a series of experiments in which the calcined pillared samples will be transferred to the NMR rotors under controlled atmosphere, in the hope that A1 coordination changes will then be noticeable. After calcination at 750 "C, all samples show significant modifications in their 27Al NMR spectrum. These are espe- cially marked for sample D750, where the global structure has already collapsed (see above), but also very significant for B750 and C750 where it has not: thus, an important trans- formation occurs in the structure of the pillars prior to the collapse of the clay layers, and possibly accelerates it.The pillars transformation seems to depend sensitively on their density in the interlayer space. In no case did we observe 27Al NMR signals reminiscent of a spinel structure. Conclusion 29Si NMR sheds light on the pillaring mechanism of All, intercalated saponites, which involves proton attack of Si-0-A1 linkages in the tetrahedral sheet followed by spe- cific inversion of AlO, tetrahedra and formation of a chemi- cal bond to the pillars, in a manner similar to other tetrahedrally substituted pillared clays such as beidellite.27Al NMR of the intercalated samples does not provide any definite evidence of the presence of species different from J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 All, in samples A, B and D, while some alumina precipi- tation may occur in sample C. 27Al NMR of the pillared samples reveals that the pillars retain their integrity after calcining at 50O0C,and suggests a reversibility of their dehydration/dehydroxylation reactions. In contrast, after calcining at 750 "C, important structural rearrangements of the pillars occurs even before the global structure collapses, but no spinel-like alumina is formed. The authors are indebted to Drs. H. Zanni and G. Herrman (Bruker Spectrospin) for facilitating access to the NMR spec-trometers, and to Dr.A. Samoson for running the DOR 4 5 6 7 8 9 10 11 12 V. H. Schmidt, in Lecture Notes, Amp2re Summer School II, Yugoslavia 1971, ed. R. Blinc, Institute R. Stefan, Ljubljana, 1971, p. 75. J. Sanz and J. M. Serratosa, J. Am. Chem. SOC., 1984,106,4790. G. Engelhardt, and D. Michel, in High-resolution Solid State NMR of Silicates and Zeolites, Wiley, Chichester, 1987, p. 150. D. Plee, F. Borg, L. Gatineau and J. J. Fripiat, J. Am. Chem. SOC.,1985, 107, 2362. J. T. Kloprogge, Ph.D Thesis, Utrecht, 1992. F. Taulelle, personal communication. J. R. Butruille, L. Michot, 0. Barres and T. J. Pinnavaia, CEAPLS Meeting on Pillared Clays, Athens, November, 1992. J. R. Butruille, Ph.D. Thesis, Michigan State University, 1992. T. J. Pinnavaia, S. D. Landau, M-S. Tzou and I. D. Johnson, J. spectra. C. Davesne conducted the experimental work on montmorillonite pillaring. 13 14 Am. Chem. SOC.,1985,107,7222. J. Sanz and J-L. Robert, Phys. Chem. Miner., 1992, 19, 39. C. A. Weiss, S. P. Altaner and R. J. Kirkpatrick, Am. Mineral., 1987,72,935. 15 G. Fu, L. F. Nazar and A. D. Bain, Chem. Muter., 1991,3,602. References 16 R. Schoonheydt, in CEAPLS Symposium on Adsorption, Separa- tion and Environmental Applications of Pillared Layered Struc- 1 S. Chevalier, R. Franck, H. Suquet, J-F. Lambert and D. Bar- tures, Antwerp, June 1993. thomeuf, J. Chem. SOC., Faraday Trans., 1994,90,667. 2 J. J. Fripiat, Catal. Today, 1988, 2, 281. 17 D. Massiot, C. Bessada, J. P. Couture and F. Taulelle, J. Magn. Reson., 1990,90,23 1. 3 S. Chevalier, Ph.D. Thesis, Paris, 1992. 18 J-F. Lambert and L. Bergaoui, unpublished results. Paper 3/04611G; Received 2nd August, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949000675

出版商:RSC    

年代:1994

数据来源: RSC