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J. CHEM. SOC. FARADAY TRANS., 1994, 90(9), 1233-1240 Profiles of Adsorption during the Oxidation of Small Organic Molecules: Oxidation of Formic Acid at Polycrystalline Pt in Acid Solutions C. Paul Wilde* and Meijie Zhang Otta wa-Carleton Chemistry Institute, University of Ottawa Campus, Department of Chemistry, University of Ottawa, 10 Marie Curie Priv., Ottawa, Ontario, Canada KIN 6N5 The electrochemical quartz crystal microbalance has been used to monitor changes in adsorption at Pt elec- trodes during the oxidation of formic acid in 0.1 mol dm-3 HCIO, solutions. This has been achieved through cyclic voltammetric, injection and open-circuit potential decay experiments where mass profiles are recorded alongside the electrochemical response. Adsorption in the H underpotential deposition (UPD) region causes the mass to increase relative to the background electrolyte whereas in the double-layer region of potential, increased coverage of strongly adsorbed intermediates has the reverse effect.Removal of these latter species can be followed from the mass response because it leads to a characteristic mass step. Subsequent to this process there is a region of potential where oxidation of formic acid occurs through consumption of adsorbed OH or PtOH and mass responses reveal that, as concentration increases, there is increased adsorption of organic residues here. The consumption of surface oxy species (OH,,, , PtOH or PtO) by formic acid also results in irreversible oxidation of the electrode surface being shifted to higher potentials with increasing formic acid concentration, since it is only at the higher potentials that the rate of the place exchange process can compete effectively with the reaction with formic acid.The mass decrease associated with removal of the surface oxide is also accelerated at higher formic acid concentrations and occurs at higher potentials. The oxidation of formic acid at Pt is a much studied reaction for two reasons. The first is interest in the area of fuel cells where formic acid might constitute the organic species to be oxidised. The second is for the elucidation of the mechanism of oxidation of small organic molecules. In one respect, formic acid is a special case since it does not necessarily require a source of oxygen (as, for example, methanol does) to produce CO,.There are several reviews of the various investigations of formic acid oxidation,'-4 but as noted else- where,4 trying to assess comparative claims in the literature is difficult. Factors such as variations in the electrode surface (e.g. roughness and orientation), the anion of the background electrolyte and its concentration and the concentration of formic acid, not to mention the different potential regimes applied to electrodes can all influence the results obtained. Thus, it is most appropriate simply to state several salient features here. First, the concept of parallel pathways for the oxidation process is generally a~cepted.~-~ Thus there is one pathway that involves a reactive surface intermediate, perhaps CO,H, adsorbed at the electrode surface, and a second that involves a chemisorbed intermediate or 'poison ', most likely CO.Both give rise to CO,, but the latter is oxi- dised only at relatively high potentials in the double-layer region on Pt and with the assistance of a source of oxygen, such as water or adsorbed OH or, at more positive poten- tials, PtOH. There is abundant spectroscopic evidence7-' ' for the presence of CO (linear, bridged or multiple site) at the electrode surface, and CHO species may also be present.12-14 The principal spectroscopic methods used to date have been various IR techniques such as electrochemically modulated IR spectroscopy (EMIRS),7-9*' ' polarisation modulation IR reflection absorption spectroscopy (PM-IRRAS)' and single potential alteration IR spectroscopy (SPAIRS),' and mass spectrometry (differential electrochemical mass spectrometry, DE M S);' ,-' surface-enhanced Raman spectroscopy (SERS)' has also recently been applied to platinum-coated gold electrodes.However, it has been suggested that different methods may be biased towards detection of different species. EMIRS for example, may detect strongly adsorbed species (CO) while DEMS can only detect adsorbates that may be desorbed from the electrode surface, and so may be more likely to identify the weakly adsorbed or reactive interme- diate.4 These factors, as well as the likelihood that the exact distribution of surface species may well be strongly dependent on the concentration of substrate, potential and the other conditions of a given experiment, should be borne in mind when considering the literature and the results to be present- ed here.Our interests are in applying the electrochemical quartz crystal microbalance (EQCM) technique to the study of elec- trocatalytic phenomena. Here, mass profiles recorded along- side electrochemical experiments add a fresh perspective on adsorption phenomena in the systems under study and have been found to be particularly valuable in locating the point where strongly adsorbed intermediates are oxidised and in monitoring the extent of oxidation of the electrode surface when this is obscured by high currents from oxidation of the organic substrate.We have applied the method previously to the oxidations of gluco~e,'~~'*and methan01'~ and to the competitive adsorption between glucose and UPD lead" and we now present results of a study of formic acid oxidation. Experimental 10 MHz AT cut quartz crystals with gold electrodes were purchased from International Crystal Manufacturing Co., Oklahoma City, OK and then platinised as described earlier.' Real surface areas of the platinum electrodes, mea- sured using integration of the H UPD charge, and the accepted conversion factor of 210 pC cm-, were typically found to be 15-20 times greater than the geometric area, which was 0.25 cm2. A saturated calomel electrode (SCE), separated from the main cell by a Luggin capillary was used as the reference electrode and a Pt wire was used as the counter (auxiliary) electrode.All stated potentials are referred to the SCE. The potentiostat used was from Oxford Elec- trodes (Abingdon, England) and current, voltage and fre- quency (mass) results were recorded on an XYY' chart recorder (Philips PM 8272 or Kipp and Zonen BD91). Control and measurement of frequency (through conver- sion of the frequency difference between working and refer- ence crystals to a voltage) was accomplished as described by Bruckenstein and Shay.21 The frequency to voltage converter has a voltage offset which allows small changes (mV) in output to be displayed when there is a high background voltage. Unless otherwise noted, this offset was not changed during a given experiment.Frequency changes were con-verted to mass as described previously,2i and for all figures mass increases up the page. Further experimental and theo- retical information on the EQCM is available in the liter- at~re.~l-~' Purified water for solutions was obtained from a Millipore Milli-Q system. Chemicals were obtained from BDH [HClO, , AnalaR; H2PtCl, , Analytical Reagent and formic acid, AnalaR (98-100%)] and Merck (H2S0,, Suprapure 96%) and were used without further purification. Stock solu- tions of 5.0 mol dm-3 (for final bulk concentrations >1 mmol dmP3) or 0.5 mol dm-3 (for final bulk concentrations below 1 mmol dmP3) formic acid were used for injection experiments. All experiments were carried out at room tem- perature, 22 & 1"C.Results and Discussion Some Generalities in the Effwt of Small Organic Fuels on Mass Responses There are several common traits in the oxidation processes of small organic molecules such as formic acid and methanol at platinum electrodes. Thus it is not surprising to find that when these processes are studied with the microbalance tech- nique some similar mass features are to be found too. These latter are presented here to provide a framework for con- sideration of the results to be discussed. (1) In general, the presence of organic adsorbates at the electrode surface causes only small differences in mass pro- files (relative to those of the background electrolyte). This is because of factors such as the small molecular weight of most organic fragments (for example CO and CHO), sub-monolayer coverages, and the fact that adsorption will involve displacement of adsorbed water and specifically adsorbed anions (typically perchlorate and sulfatelhydrogen sulfate).Furthermore, one CO, for example, may occupy one, two or three surface sites. In fact, for Pt electrodes in HClO, , it has been found that coverage of the electrode surface with poisoning adsorbates in the double-layer region of potential (whether they be derived from gluco~e,'~ methanolig or formic acid) leads to a mass decrease. It should also be noted here that because of the factors mentioned above, together with the difficulty in obtaining an exact indication of the nature of distribution of species adsorbed at the surface as a function of potential, the extraction of coverage values for adsorbates from mass data is exceptionally difficult.(2). The removal (by oxidation) of strongly adsorbed species (which is often associated with a sharp peak in the voltammogram) leads to a transition between a surface that is largely covered with organic adsorbates and one that is largely free from such adsorption (this is less true at concen- trations > lop2mol dmP3). As a consequence, a mass step that corresponds to a mass increase (see comment above) is often observed to accompany the removal process. However, the likelihood of adsorption subsequent to the removal of the strongly adsorbed intermediates increases with the concentra- tion of the organic reactant.Thus the mass step diminishes in size as concentration increase^.'^ (3) For the oxidations of 0.1 mol dm-3 glucose in alkaline media,'* of 0.1 mol dm-3 methanol in perchloric acid solu- tions,lg and of 0.1 mol dm-3 formic acid (to be presented here) the mass responses reveal a significant positive shift in J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 the point at which irreversible oxidation of the electrode surface begins. (The large current due to organic oxidation prevents detection of this phenomenon in the voltam-mogram.) In both background electrolyte and when organic fuels are present, there is a region of constant mass before the onset of irreversible surface oxidation.The mass increases again only when the place exchange process [insertion of OH (and 0)into the Pt latti~e~~.~'] starts, because the initial stages of oxidation of the surface, where PtOH is produced from adsorbed water, do not lead to a mass change. The point where mass begins to increase again after the double- layer region of potential, and its variation with increasing amounts of organic fuel in the electrolyte are thus easily iden- tified. A related item of interest is that the rate of increase of mass (and hence oxide growth) is more rapid when shifted towards more positive potentials, presumably because of the increased field at the higher onset potential. We will now illustrate and elaborate upon some of the above comments by examining the mass responses resulting from the oxidation of formic acid at various concentrations in 0.1 mol dm-3 HClO, .Oxidation of Formic Acid at Very Low Concentrations, 0.5 mmol dm-3 Fig. 1 shows a voltammogram and accompanying mass response, recorded after the potential profile in the inset, for 0.5 mmol dmV3 formic acid. Essentially there is a broad oxi- dation of HC02H across almost the entire potential range. This oxidation current first increases above the background at around -0.1 V. Current densities are low, and the roughly constant oxidation is disturbed only by the sharp peak in the double-layer region of potential on the anodic scan. The first 40 20 0 -20 5.: -40 I 1 -0.2 0.2 0.6 1.o E/V vs.SCE Fig. 1 Cyclic voltammogram and mass response for the oxidation of 0.5 mmol dmP3 formic acid in 0.1 mol dm-3 HC10,. Results for the background electrolyte alone are also shown (dotted line). El = 1.15 V,t, = 120 s, E, = -0.25 V and tpds= 60s. The scan rate was 5 mV s-and the electrode area 3.82 an2.Mass responses are shown exactly as recorded. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 stages (to -0.1 V) of the mass response accompanying the anodic scan are seemingly unaffected by the presence of formic acid. (A discussion of the mass response in 0.1 mol dm-3 HClO, alone has been presented elsewhere.I7) This is not unexpected, given that the value of E, was -0.25 V where H coverage is nearly unity. Capon and Parsons6 have shown that the reaction between Hadsand HC0,H (bulk) is slow in such a circumstance.From -0.1 V on, however, the mass response is flat and lower than in the background elec- trolyte over a range of nearly 400 mV (indicated as region A) until the step (feature B) that accompanies the sharp peak. Here then are two of the standard features described earlier. First, the steady region of mass reflects what is probably a constant (small) coverage of strongly adsorbed species, pre- sumably developed in the earliest stages of the anodic scan and then the step corresponds to their removal. The electrode surface thereafter seems to be quite similar to that seen in the background electrolyte since there is little difference in mass. (Note that the two mass traces are shown exactly as recorded and it is most likely that in fact the masses should be equal at the anodic scan limit.Some evidence for this is presented later, in the discussion of mass transients. The difference in the figure is probably a result of a small drift in the frequency between the recording of the scans.) When present at this low level, formic acid has no influence on the oxidition-reduction of the electrode surface and so the next point of interest follows oxide removal on the negative-going scan. Here, there is another flat region, yet the mass is larger than in the corre- sponding section of the positive-going scan. This difference arises because the double-layer region is approached from opposite directions.On the anodic scan the initial sections involve some residue formation (removed in the sharp peak) before the double-layer region is encountered, whereas on the cathodic scan the surface in the double layer region is ‘clean’ since it is reached after oxide removal. These data indicate that, under the conditions of this experiment, the formation of strongly adsorbed species occurs principally in the region of weakly adsorbed H. Fig. 2 supports this assumption. Here, 1 I 1 I I I 1 I 0.0 0.4 0.8 EfV vs.SCE Fig. 2 Details as for Fig. 1 except that E, was -0.10 and 0.10 V for curves 1 and 2, respectively. Electrode area 4.88 cm’. The vertical arrow labelled 2 indicates the starting point for the mass for curve 2. the potential E, was increased to -0.1 V and 0.1 V as com- pared to -0.25 V in Fig.1. For E, =-0.1 V, the mass profile mirrors that of Fig. 1 and the sharp current peak is seen (though there is less charge involved). For E, =0.1 V, however, the peak is negligible (the current is constant almost throughout the double-layer region) and the mass on the anodic scan actually follows the path of the signal from the cathodic scan until it merges with the mass profile from the first anodic scan at 0.4 V at the beginning of the plateau region. Thus under the conditions of these experiments, we may conclude that a small amount of adsorption occurs prin- cipally at the extreme negative limits of the scan, and there is little further adsorption in the double-layer region of poten- tial.This is not unexpected, given the small bulk concentra- tion used. The general picture of a constant small oxidation current across virtually the whole potential range, together with the sharp peak, seems to fit well with the accepted view that the oxidation of formic acid occurs by parallel routes with a reactive intermediate and a strongly adsorbed interme- diate. The differences in mass response in the double-layer region of potential also illustrate the fact that increased coverage here causes the mass to decrease and a comparison of the mass responses can thus give a qualitative picture of the variation of coverage between the two scan directions. We do not believe that changes in surface roughness (with possible increased solvent entrapment in pores) play any role in causing these differences in the mass response in the double-layer region of potential. This is for two reasons.First, the surfaces are already rough (see Experimental section) and significant changes in roughness upon cycling are thus not to be expected. Secondly, the response in back- ground electrolyte (Fig. 1, dotted line) shows no difference (within experimental error) between the mass responses accompanying the anodic-going and cathodic-going halves of the cycle. Changes in roughness upon cycling through oxidation-reduction of the electrode would not result in the mass response forming a closed loop as it does here. As a final point, identical experiments to those of Fig.1 and 2 but with methanol” showed that the formation of strongly bound intermediates began at a more positive potential (than for formic acid) and continued through most of the double- layer region leading to a continued decrease in mass, rather than the flat response seen here. Thus subtle differences between the behaviour of different reactants can be discerned from the mass responses. Cyclic Voltammograms for the Oxidation of 9 mmol dm-3 Formic Acid When the bulk formic acid concentration is increased to 9 mmol dm-3, changes in both the voltammetry and mass response become apparent. Currents are larger and an anodic scan is dominated by the sharp peak at 0.4 V with a steady increase in current prior to this and a shoulder thereafter. [Note: The current before the sharp peak is often apparent as another peak, frequently referred to as the first anodic peak.6 Under the conditions used here, namely slow scan rates and 0.1 mol dmP3 HClO, as electrolyte, this first anodic peak was not observed (cf:ref.15, Fig. 7). However, at faster scan rates in 0.5 mol dm-3 H,SO, it was present.] On the cathodic scan, there is a reactivation peak which is followed by a dip due to surface oxide reduction and then a current plateau before the final decline in current begins at 0.1 V. Fig. 3 also shows the second cycle recorded directly after the first. The only significant difference between the voltammograms is the suppression of the current in the region from -0.25 V to -0.1 V. This is a result of the higher coverage of strongly adsorbed species formed in the latter stages of the first cycle.1236 0.2 0.1 0.0 5-0.1 --.\ 0.0 0.4 0.8 E/V vs. SCE Fig. 3 Cyclic voltammogram and mass response (upper mass trace) for the oxidation of 9 mmol dm-3 formic acid in 0.1 mol dm-3 HClO,. Results were recorded after a potential profile like that of Fig. 1. The second cycle was recorded directly after the first with no pause. The electrode area was 4.88 cm2 and the scan rate was 5 mV s-'. The background mass response is shown only for comparison purposes and is displaced downwards for purposes of clarity. The vertical dotted line is drawn to illustrate the slight shift in the point where irreversible surface oxidation begins when formic acid is present.Point C indicates where the two anodic-going mass responses coincide and point D illustrates the mass plateau that follows the step. Mass Responses for 9 mmol dm-3 Formic Acid, Variation of Electrode Mass with Coverage in the H UPD Region and the Double-layer Region The mass responses for the two cycles with formic acid, in the middle of Fig. 3, are in excellent agreement, except during the initial stages of each of the anodic halves of the two cycles. The reasons for this difference will now be addressed. The state of the surface at a potential of -0.25 V is affected by the potential perturbation applied to the electrode before it reaches -0.25 V. In the first cycle there is a step from 1.15 V and a pause of 60 s at -0.25 V, where some adsorption will take place.This is likely to be small, as discussed above. Further adsorption can take place once the potential scan begins and adsorbed H is removed from the surface. In con- trast, on the second cycle the anodic scan is preceded by a cathodic scan where, once the surface oxide is removed, there is a considerable time for coverage of adsorbates to develop. Comparison of current and mass in this section of the cathodic scan is interesting since in a general sense the two signals are similar. Thus the region of constant current from 0.4 V downwards is accompanied by a region of almost con- stant mass and as the current decreases more sharply so does the mass, at least for a short time.Finally from about 0.0 V the mass increases, until it reaches -0.25 V. It is clear from a comparison of the voltammetry and mass that a higher coverage of adsorbates leads (in the H UPD region) to an J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 increase in mass, whereas the reverse is true in the double- layer region of potential. However, the fact that responses for the two cycles merge at ca. 0.15 V implies that at this point (C) there is a comparable coverage (the voltammetry is quite similar too) which is reached by different routes on the two scans. It is more difficult to explain why the mass responses take the shape they do and why an increased coverage of adsorbates leads to a mass increase in the H UPD region and a mass decrease in the double-layer region.Undoubtedly some of the factors described in the introduction play a part. For example, the coverage of water and specifically adsorbed anions is potential dependent and thus one might not neces- sarily expect that the occupation of an equal number of sites by a given adsorbate results in the same net mass change at two different potentials. We must also recognize the roles of time and formic acid concentration as additional influences because the rate of adsorption will vary with both of these factors. Thus the scan rate and the bulk formic acid concen- tration will influence the mass response. Further information on this point can be obtained from mass-transient experi- ments, but before presenting those results the remainder of the cycles of Fig.3 will be discussed. Mass Responses for Oxidation of 9 mmol dm-3 Formic Acid, Removal of Strongly Adsorbed Species and Electrode Surface Oxidation Once the sharp peak is reached in the voltammogram, the mass step is seen. There is then a plateau region (D) before the mass increases again. An interesting question is whether or not there is any adsorption in this plateau region. Certain- ly we know that prior to the mass step there is a significant coverage of strongly adsorbed intermediates, but what is the situation thereafter? If one chooses a particular potential in the mass plateau region, say 0.5 V, and then measures the mass difference between this point and the mass at the poten- tial of scan reversal, the difference is larger than when a similar measurement is made for the background electrolyte alone.If one assumes that the mass at scan reversal is the same in both cases (i.e. that there is a similar coverage of oxide and that there is little adsorption of organic residues; both of these points are substantiated by later results) then this implies a small amount of adsorption in the plateau region around 0.5 V. This makes the electrode mass lower than in 0.1 mol dmP3 HClO,. It is interesting to note that the mass plateau region corresponds to a broad shoulder in the voltammogram. It may be that (some of) the oxidation here takes place with the assistance of adsorbed OH or PtOH at the electrode surface. The likelihood of adsorption in this region is also indicated by the small shift in the onset of the final mass increase of the anodic scan (dotted line).This implies some blockage or shifting of the irreversible stages of surface oxidation. There is little apparent effect of the formic acid on the reduction of the surface oxide, but it is interesting to note that the reactivation peak on the cathodic half of the voltammogram occurs before any noticeable decrease in mass implying that no significant reduction of the oxidised surface is necessary to allow the oxidation of formic acid to begin again. Mass Transient Experiments One simple means of examining adsorption processes at the EQCM is to add a small amount of the adsorbate to the electrolyte and follow the frequency (mass) signal with time at constant potential. In general, changes in solution viscosity and density during such experiments are too small to influ- J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Fig. 4 Mass transients resulting from the addition of formic acid (sufficient to give a bulk concentration of 9 mmol dm-3) to the back- ground electrolyte at different constant potentials. Each injection was performed at a clean electrode surface. Electrode area 4.30 an2. Responses are displaced for purposes of clarity and are not intended to represent relative mass values. Constant potentials used were (a) -0.15, (b)0.0, (c) 0.15, (6)0.30, (e) 0.55 and (f) 1.0 V. The arrows indicate when the formic acid was added to the electrolyte, which was stirred with nitrogen so that the time of mixing is just a few seconds.ence the mass response, although this was found not to be the case when concentrated glucose solutions were generated by injection.20 Thus it is possible to see whether adsorption leads to a simple mass increase or decrease relative to the background electrolyte. The progression of the mass to a steady state or pseudo-steady state can also be followed. For CO adsorption, for example, initial adsorption at Pt leads to a mass decrease but as coverage develops the mass change reverses direction and finally a net increase is seen.28 These types of experiment thus provide further insight into adsorp- tion processes at constant potential and to the relative posi- tions of mass responses accompanying vol tammograms for background electrolyte containing different amounts of formic acid from zero upwards. Results of some mass transient experiments are shown in Fig.4. Additions of formic acid were chosen to provide a final bulk concentration of 9 mmol dm-3. Each experiment was carried out at a clean electrode. The results suggest the following general conclusions. First, at the oxidised electrode surface (1.0 V) there does not appear to be any significant adsorption. This was found to be true for concentrations up to 1.0 mol dm-3. Secondly, in the double-layer region of potential (0.15,0.30 and 0.55 V) there is adsorption such that the mass of the electrode is less in the presence of formic acid.Both these conclusions agree with those derived from mass responses that accompanied the voltammetry. (The potential dependence of the mass change also reinforces the conclusion that there is no effect of the small change in solution viscosity and density on the mass response.) Finally, at 0.0 V and -0.15 V a small increase is seen in the mass of the electrode after addition of formic acid. This too is what would be expected from the data in Fig. 3, although it should be appre- ciated that the mass response accompanying the CV is depen- dent upon processes occurring at lower potentials. The result at 0.0 V also shows the effect of time, since the initial mass change is a decrease followed by a steady increase. Again this is most likely to be a result of the competition between anions and water and the adsorbate and perhaps also of the changing nature of the adsorbate with coverage (e.g. CO may [HCO,H]/mmol dm-3 II' 1 I I B I1 I 1 I I 0.3 0.5 0.7 0.9 1.I E/V vs.SCE Fig. 5 A, Variation of the potential (E,) of the anodic (m) and cathodic (reactivation) peaks (0)with formic acid concentration. Scan rate was 5 mV s-' and all voltammograms were recorded after a potential profile like that shown in Fig. 1. B, Influence of formic acid concentration on mass response. (a) 20 mmol dm-3, (b)40 mmol dm-3, (c) 60mmol dm-3 and (d) 0.1 mol dm-'. The concentration of formic acid was increased in steps and then each response recorded after the potential profile shown in Fig.1 at 5 mV s-'. Traces are drawn aligned so that the masses coincide at the upper potential limit. Electrode area 4.55 cm2. shift from being multiply bound to being singly bound as coverage increases). The combined data provide a useful illus- tration of the fact that mass changes upon adsorption can be either positive or negative. Influence of Increasing Amounts of Formic Acid on the Mass Profile An increase in the amount of formic acid present in the elec- trolyte does not cause significant changes in the cyclic volt- ammetric response, but the following general trends are observed (when all voltammograms are recorded after a potential profile like that seen in Fig. 1). First the anodic peak (identified with the second anodic peak of Capon and Parsons6) that is associated with the mass step (and with removal of strongly adsorbed species) is shifted steadily in the positive direction from 0.42 V at 20 mmol dm-3 to 0.70 V at 0.1 mol dm-3.The shoulder following the peak also grad- ually disappears until there is, at 0.1 mol dm-3 formic acid, an abrupt drop in current after the peak. At the same time the reactivation peak (where formic acid is oxidised on fresh sites generated by some removal of the surface oxide) becomes much sharper. It also shows a positive shift, but to a lesser extent. These data are summarised in Fig. 5A. Needless to say, current densities also increase. The evolution of the mass response is more interesting. Fig. 5B shows part of the mass responses (restricted to potentials in the double-layer region and above) for a series of experiments on the same electrode.There are three general effects. First, the mass step shifts in a positive direction, just as the anodic current peak does. It also diminishes in size and the subsequent mass plateau shortens. Second, the final mass increase of the scan is also shifted in the positive direction. Finally, the rate of removal of surface oxide on the cathodic scan is increased and thus the mass decreases more sharply than in the back- ground electrolyte and at a more positive potential. This behaviour is very similar to that reported earlier for methanol’ and leads to several conclusions. (1) The removal of strongly adsorbed species is shifted towards more positive potentials as the bulk concentration increases.This results in a similar shift of both mass step and current peak. This may occur because there are less oxy species (OH,,, or PtOH) at the surface to assist in oxidation of CO as a result of more extensive occupation of the surface by the reactive intermediate and possibly by CO itself. (2) There is probably more adsorption subsequent to the mass step as concentration increases, but there comes a point where the mass plateau region has almost disappeared (and so has the shoulder in the voltammogram after the peak, as noted above). The gradual shift in the potential at which oxi- dation of strongly adsorbing species occurs leads to a situ- ation where the removal process is immediately followed by irreversible surface oxidation.This effectively ‘switches off’ the formic acid oxidation process, although it is well known that the oxidation begins again at more positive potentials on the fully oxidised surface. (3) The more rapid reduction of the oxidised surface, illus- trated by the rapid decline of the mass earlier in the cathodic scan (and the positive shift in the reactivation peak) is most likely to be associated with an increased reaction between adsorbates derived from formic acid and PtOH as concentra- tion increases. This point is dealt with further in the next section. A complete cycle for 0.1 mol dm-3 formic acid with both voltammetric and mass responses can be seen in Fig.6, com-pared to the background responses. This illustrates more clearly the changes that occur as the formic acid concentra- tion is increased, particularly the almost complete absence of the mass step on the anodic scan and the significant shifts in the potentials of oxide formation and removal. Thus the final increase in mass which corresponds to the beginning of irre- versible surface oxidation is shifted to 0.79 V compared to 0.57 V in the background electrolyte. It is more rapid at first and then slows to a rate similar to the background response. The faster removal of the surface oxide is also seen to coin- cide clearly with the sharp reactivation on the negative-going scan. Two further small points should be made about the response at this concentration.In the double-layer region of potential the mass is flat, and there is no difference between the anodic and cathodic sections. It is likely that at this con- centration the steady-state coverage of adsorbates is reached rapidly (clearly this is the case after oxide removal) and is the same on both halves of the cycle. As before, the mass increases as the potential proceeds into the H UPD region (on a cathodic scan) since a higher coverage of adsorbates in this region leads to a mass increase. The slight discrepancy between the beginning and end of the mass loop may be accounted for by differing coverages. Finally the relative posi- tions of the two mass responses are again substantiated by injection (mass transient) experiments which may be sum-J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 4.0 3.0 2.0 41.o .-.\ 0.0 Fig. 6 Cyclic voltammogram and mass response for 0.1 rnol dm-3 formic acid in 0.1 mol dm-3 HCIO,. Scan rate 5 mV s-l, electrode area 4.55 cm2. The background mass response (dotted line) is pre-sented for purposes of comparison and is drawn so that masses coin- cide at the point of scan reversal. The vertical arrow between points F and G represents the mass change when 0.1 mol dm-3 formic acid is added to the electrolyte with the potential held at 0.7 V. See text for details. marised thus. In the H UPD region the mass of the electrode is larger than in the background. (The observed increase is in fact larger than that seen for the same potential with 9 mmol dm-3 formic acid.) In the double-layer region the mass of the electrode is less than in the background electrolyte and finally, on the oxidised electrode surface any mass difference is at the limit of experimental significance.Role of Surface Oxy-species in the Oxidation Reaction, Experiments with 0.1 mol dm-3 Formic Acid Injection and open-circuit potential decay experiments, where mass is recorded in parallel, provide some interesting further information regarding the interaction between formic acid and the electrode surface when it is at least partially oxidised. The injection experiment was carried out with the electrode held at 0.7 V (after a step up from 0.15 V) in background electrolyte. Formic acid is then added to give a final concen- tration of 0.1 mol dm-3 while the potential is maintained at 0.7 V.This causes a drop in mass which is larger than could be accounted for by simple reduction of the (partially) oxi- dised surface alone. The mass decreases because the partial coverage of oxide is removed and then there is adsorption of formic acid or intermediates. This process can be represented in Fig. 6 by a shift from a point at 0.7 V on the background response (F) to a corresponding point on the response with 0.10 mol dm-3 formic acid (G).At this potential, electro- chemical oxidation (turnover of PtOH to OHPt) does not appear to be able to compete effectively with removal of OH or PtOH by reaction with formic acid or intermediates. However, if the potential is increased to 1.0 V and the same injection made, there is no change in mass (note that this is J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 effectively the experiment shown in Fig. 4 except that the bulk concentration then was 9 mmol dm-3). Here, the rate of electrochemical regeneration of any oxide removed in a chemical reaction is sufficient to prevent a noticeable decrease in mass. It is also noted that at this potential there is no difference between the mass responses as drawn in Fig. 6. That the chemical reaction between the oxidised surface and formic acid is possible is well known2 and the results of Fig. 7 illustrate this. The potential was taken to 1.0 V (in back- ground electrolyte) and the circuit opened.Once a stable potential was attained, formic acid was added to give a chosen final concentration and the potential and mass were followed with time. Results are shown for two concentrations 0.5 and 100 mmol dmP3. Addition of formic acid in both cases leads to a shift in the potential to the region of H UPD, as expected, and to a decrease in mass. Two further points emerge. The rate of decrease is slower for the lowest concen- tration as might be expected, but there is also a larger mass decrease for the lower concentration. Part of this decrease is accounted for by the removal of the oxide, but this should be equal in both cases, The remainder represents the difference between the stable relative masses of the electrode surface in the presence of differing amounts of formic acid.It is there- fore suggested that at the potentials of interest (in the H UPD region) there is a larger coverage of adsorbates devel- oped from 0.1 mol dm-3 formic acid and hence the mass is larger and so the change in mass seen after the potential- decay experiment is smaller for 0.1 mol dmP3 formic acid. Again one can carry out an injection experiment to investi- gate this and indeed at a constant potential of -0.15 V titra-I \ I0.20 v \ \ Fig. 7 Mass (upper) and potential (lower) transients observed upon addition of sufficient formic acid to yield a final bulk concentration of 0.5 mmol dm-3 (dashed line) and 0.1 mol dm-3 (solid line) in 0.1 rnol dm-HCIO,. The electrode was held first at 1.0 V and then the circuit was opened.The starting potential was 0.8 V in both cases, and the final potentials were -0.11 V (0.5 mmol dm-3) and -0.18 V (0.1 rnol dmP3). The arrow indicates when formic acid was added. Electrode area 3.50 cm2. tion of the formic acid concentration to higher amounts leads to successive increases in mass. Conclusions Despite the most obvious drawback of the technique, namely the inability to identify adsorbed species present at the elec- trode surface, significant information can be obtained through the careful application of the EQCM to electro- catalytic reactions. Interactions at the electrode surface are clearly complex and the fact that the net mass change upon adsorption depends upon the species displaced from the surface as well as those adsorbed is well illustrated.Thus in the H UPD region of potential, increased coverage of the electrode surface by adsorbates derived from the reaction of formic acid causes an increase in the mass of the electrode. In contrast, in the double-layer region of potential the reverse is true and differences in the mass in this region on the two halves of a scan can reveal the differing extents of adsorption. The mass step is found to be a feature common to several processes and is useful in locating the point where oxidative removal of strongly bound adsorbates occurs. The sub- sequent mass plateau region is also informative since it can be correlated with a shoulder on the voltammogram and it is most likely that oxidation here involves consumption of surface OH, either adsorbed or as PtOH.The evolution of these two features with concentration shows a shift of the step to higher potentials, and increased adsorption following the step (so that it becomes smaller in size). Perhaps the most useful data are those furnished from considering the oxidation-reduction of the electrode surface, and how it is influenced by varying amounts of formic acid. Mass results accompanying cyclic voltammetric and injection experiments data show that higher concentrations of formic acid appear to consume surface oxy species (perhaps PtOH) at a rapid rate in the mass plateau region so that irreversible surface oxidation is shifted towards more positive potentials until the rate of the place-exchange process will have increased sum- ciently to compete with consumption by formic acid (it is also possible that there is some blockage of surface sites by stongly adsorbed species).The rate of growth of the oxide (as represented by the rate of increase of mass) is then faster than usual when it does occur because of the higher field. In addi- tion, removal of the oxide occurs at a more positive potential. This is also a result of the swift consumption of PtOH and PtO by formic acid. We thank the Natural Sciences and Engineering Research Council of Canada for financial support (a Research Grant) and M.Z. wishes to thank the Government of Ontario for the award of an Ontario Graduate Scholarship and the Cana- dian Local Section of the Electrochemical Society for the Student Award (1993).References M. W. Breiter, Electrochemical Processes in Fuel Cells, Springer-Verlag, New York, 1969, 157. W. Vielstich, Fuel Cells, Wiley, New York, 1970, p. 76. A. Capon and R. Parsons, J. Electroanal. Chem., 1973,44, I. R. Parsons and T. VanderNoot, J. Electroanal. Chem., 1988,257, 9. A. Capon and R. Parsons, J. Electroanal. Chem., 1973,44,239. A. Capon and R. Parsons, J. Electroanal. Chem., 1973,45205. B. Beden, A. Bewick and C. Lamy, J. Electtoanal. Chem., 1983, 148, 147. B. Beden, A. Bewick and C. Lamy, J. Electroanal. Chem., 1983, 150, 505. S. G. Sun, J. Clavilier and A. Bewick, J. Electroanal. Chem., 1988, 240, 147. 1240 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 10 11 12 13 14 15 16 17 18 K. Kunimatsu, J. Electroanal. Chem., 1986,213, 149. K. Kunimatsu and H. Kita, J. Electroanal. Chem., 1987, 218, 155. J. Willsau and J. Heitbaum, Electrochim. Acta, 1986,31,943. 0. Wolter, J. Willsau and J. Heitbaum, J. Electrochem. SOC., 1985,132,1635. N. A. Anastasijevic, H. Baltruschat and J. Heitbaum, J. Electro-anal. Chem., 1989,272,89. D. S. Corrigan and M. J. Weaver, J. Electroanal. Chem., 1988, 241, 143. Y. Zhang and M. J. Weaver, Langmuir, 1993,9, 1397. C. P. Wilde and M. Zhang, J. Electroanal. Chem., 1992,340,241. C. P. Wilde and M. Zhang, J. Chem. SOC., Faraday Trans., 1993, 20 21 22 23 24 25 26 27 28 C. P. Wilde and M. Zhang, Electrochim. Acta, 1993,38,2725. S. Bruckenstein and M. Shay, Electrochim. Acta, 1985,30, 1295. D. A. Buttry, in Electroanalytical Chemistry, ed. A. J. Bard. Marcel Dekker, New York, 1990, vol. XVII, p. 1. R. Schumacher, Angew. Chem., Zntl. Ed. Engf., 1990,29,329. M. D. Ward and D. A. Buttry, Science, 190,249,1000. D. A. Buttry and M. D. Ward, Chem. Rev., 1992,92, 1355. H. Angerstein-Kozlowska, B. E. Conway and W. B. A. Sharp, J. Electroanal. Chem., 1973,43, 9. B. E. Conway, Prog. Sur$ Sci., 1984, 16, 113. S. Bruckenstein and C. P. Wilde, Ext. Abs. 116, 198th A.C.S. National Meeting, Miami Beach, FL, 1989. 19 89, 385. C. P. Wilde and M. Zhang, Electrochim. Acta, 1994,39, 347. Paper 3104942F; Received 16th August, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949001233

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(9), 1241-1244 Influence of Polymer Structure on the Electrochemistry of Phenothiazine Dyes incorporated into Nafion Films Swamidoss A. John and Ramasamy Ramaraj" School of Chemistry, Madurai Kamaraj University, Madurai-625 021, India Electrochemical properties of phenothiazine dyes, thionine (TH+) and methylene blue (MBf), incorporated into Nafion films have been studied at different acid concentrations. The TH+ and MB+ dyes adsorbed by Nafion films gave an unusual redox wave at more positive potentials in addition to the usual redox wave at less positive potentials in 0.05 mol dm-3 H,SO,. Upon continuous cycling for 45 min the unusual redox wave disappeared. The two redox waves observed at 0.05 mol dm-3 H,SO, were accounted for by the distribution of electroactive species in different domains of the Nafion polymer film, i.e.the lower void volume of the interfacial region and the higher void volume of the ionic cluster region. The same dyes incorporated into Nafion films in 0.5 mol dm-3 H2S04 gave only the usual redox wave, which was accounted for by the molecules residing only in the ionic cluster region of the swelled Nafion film. The same electrochemical behaviour was also observed in Na+- and H+-Nafion film. The effects of different concentrations of dyes in the Nafion films were also studied in order to understand the influence of the polymer film structure on the electrochemical properties of phenothiazine dyes. Electrodes modified with the perfluorosulfonate membrane, with low charge (e.g.ferrocene) gave an additional redox Nafion, have received much attention, owing to its outstand- wave. However, highly charged molecules (e.g. [Ru(bpy): '3)ing chemical and thermal stability.' An electrode surface did not give any additional oxidation or reduction waves. '' modified with Nafion film containing electroactive species There are only a few reports in the literature on the elec- exhibits interesting electr~catalytic~.~ and photo~hemical~,~ trochemical properties of dye molecules incorporated into properties. In the last decade, much effort has been devoted to elucidate the mechanisms of diffusional charge transport in Nafion polymer films at electrode surfaces.'*' Nafion membrane provides a unique solid matrix contain- ing a fluorocarbon backbone with pendant chains terminated with sulfonate head groups which are responsible for its ion- exchange and swelling proper tie^.^,^ The cluster network model of Gierke and Hsu' described the ionic regions of Nafion film as spherical with interconnecting channels. Later Yeager and Steck" proposed a three-phase structure, i.e., a hydrophobic fluorocarbon phase, a -SO; ion cluster region and an interfacial region between these two (Fig.1). The effects of the water uptake characteristics of the Nafion film,'' curing humidity," and variations in the properties of Nafion due to different preparation condition^'^ have also been subjects of recent interest. The first direct evidence for the existence of multiple-phase structure in a Nafion film was reported by R~binstein'~ using cyclic voltammetric experi- ments, This was later supported by further evidence from Vining and Meyer" and Harth et a/.'' that uncharged mole- cules gave an additional oxidation wave, whereas molecules Fig.1 Schematic structure of Nafion polymer film: H, hydrophobic fluorocarbon phase; C, hydrophilic -SO; ionic cluster region; I, interfacial region; g,SO; group of the polymer Nafion-film-coated electrodes."-19 Studies were mainly carried out to investigate the charge-transport properties of the dyes and also to elucidate the dimerization kinetics of the dye molecules in Nafion film^.'^.'^ Even though electro-chemical behaviour such as the transport properties and absorption spectral characteristics of thionine (TH +) and methylene blue (MB') at Nafion-coated electrodes has been investigated, the influence of Nafion-film structure on these dyes has not yet been studied.These dye-modified electrodes found applications in solar energy conversion, and electro- chromic and photochromic devices.20-22 In the present paper, we report the observation of an unusual redox wave for thionine (TH') and methylene blue (MB') dyes incorpor- ated into Nafion-film-coated electrodes in 0.05 mol dm-3 H2S0,. Studies were also carried out using Na+- and H+- Nafion films and different dye concentrations in the Nafion film. Experimental Materials A 5% solution of Nafion 117 (Aldrich, 1100 EW) was diluted to 1-2% with methanol before use.Thionine and methylene blue (Aldrich) were purified by chromatographic methods.23 All other chemicals were of analytical grade and used as received. Apparatus and Methods A three-electrode cell was used with 1 cm2 platinum plates as working and counter electrodes and a saturated calomel elec- trode (SCE) as the reference. Cyclic voltammograms were run using an (EG&G) PAR 273A potentiostatlgalvanostat equipped with an RE0151 recorder. The surface coverage (r)was determined using a coulometric method by measuring the charge required in a potential-step experiment to reduce the dye molecules q~antitatively.~~ The diffusion coefficients were determined by chronoamperometry.25 Nafion films were prepared by casting a known volume of 1% or 2% Nafion solution onto a 1 cm2 platinum plate and the solvent was evaporated from the surface of the electrode at room temperature for 10 min.Then, the Nafion-coated electrode was washed with distilled water and stored in dis- tilled water for 30 min. The thickness of the Nafion film was calculated from the amount of Nafion present in the solution and the density of Nafion (2 g ~rn-~).’~,” It has been reported” that the Nafion film thickness did not change sig- nificantly when the film was immersed in 0.2 mol dm-3 Na,SO, for 2 h. The Na+-Nafion film was prepared by soaking the Nafion-coated electrode in dilute NaOH solution and the H+-Nafion film by soaking in dilute H2S04.26*27 The dye-incorporated Nafion-film-coated electrodes were prepared by dipping the Nafion-coated electrode in a solu- tion containing a known concentration of dye.The solutions were purged with nitrogen for 30 min before each experiment to exclude dissolved oxygen. Results and Discussion Recently, the influence of Nafion-film structure with hydro- philic and hydrophobic environments on molecules of low charge was reported.16 Such an effect was not noticed in the case of phenothiazine dyes under the experimental conditions employed in previous studies.’ 7-19 Earlier studies of TH’ and MB+ incorporated into Nafion films in 0.5 mol dmd3 H2S04 have shown a well defined reversible redox wave in the potential region of +0.4-0.0 V us. SCE.17-” Cyclic voltammograms recorded for the Nafion-coated electrode dipped into a mixture containing TH+ and 0.05 mol dm-3 H2S04 are shown in Fig.2(a). In the first cycle, TH’ in the Nafion film gave two well defined redox waves at 0.16 and 0.39 V us. SCE. The latter wave is unusual for TH’ in Nafion films.” Upon continuous cycling, the wave at 0.39 V us. SCE decayed and that at 0.16 V us. SCE increased [Fig. 2(a)]. When the Nafion-coated electrode was dipped into a mixture containing TH’ and 0.5 mol dm-3 H2S04, the redox wave at 0.39 V us. SCE disappeared during the first cycle and a wave at 0.18 V us. SCE appeared, as shown in Fig. 2(c). Very similar behaviour was also observed for MB+ using Nafion-coated electrodes [Fig. 2(b) and (43.The cyclic voltammogram recorded for a Nafion-coated electrode dipped into a mixture containing MB+ and 0.05 mol dm-3 H2S04 showed two redox waves at 0.14 and 0.40 V us.SCE [Fig. 2(b)]. When 0.5 mol dm-3 H2S04 was used as support- ing electrolyte, only the usual redox wave was observed at 0.18 V us. SCE [Fig. 2(4]. I I J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 In the case of ferrocene, the cyclic voltammograms showed an unusual wave, in addition to the usual redox wave.l4.l6 It was suggested that ferrocene (Fc) and ferrocenium (Fc’) (a neutral molecule and an ion of low charge, respectively) occupy different domains, such as the sulfonate ionic cluster region and an interfacial region existing between the hydro- phobic fluorocarbon domain and the ionic cluster region.However, in the present study the TH’ and MB’ molecules possess one positive charge and upon two-electron reduction these dye molecules form leuco-dyes with one positive charge [eqn. (1)and (2)].’79’9>28 TH+ + 2H’ + 2e--,TH; (1) MB+ + 2H+ + 2e-+MBH; (2) These dye molecules are hydrophobic in nature and upon adsorption by Nafion film in the presence of 0.05 mol dmd3 H2S04, might occupy both the ionic cluster and interfacial regions, as in the case of ferrocene molecules,’6 to give two redox waves, as inferred in Fig. 2(a) and (b). Furthermore, upon continuous cycling, the dye molecules present in the interfacial region might be transported to the ionic cluster region, owing to the involvement of H+ ions in the redox reaction [eqn.(1) and (2)J Hence the unusual redox wave decreases and the usual redox wave increases. The unusual redox waves obtained for TH+ and MB+ are assigned to molecules adsorbed by the interfacial region of the Nafion film. Upon increasing the acid concentration, the dye mol- ecules are adsorbed by the well solvated ionic cluster region of the Nafion film and the unusual redox waves are no longer observed. The cyclic voltammograms obtained for MB’ incorpor-ated into Na+- and H+-Nafion-coated electrodes in 0.05 mol dmd3H2S04 exhibit two redox waves as shown in Fig. 3(a) and 3(b). The corresponding electrochemical data are given in Table 1. The amounts of electroactive species for the two redox reactions were obtained coulometrically in the first cycle by stepping the potential from +0.7 to +0.3 V and from +0.3 to -0.1 V us.SCE. The amount of electroactive species in 0.5 mol dm-3 H2S0, is obtained by stepping the potential from +0.7 to -0.1 V us. SCE. The corresponding electrochemical data are given in Table 2. The amount of electroactive TH’ and MB’ species in the H+-Nafion film is higher than that in the Na+-Nafion film. Note that it has already been reported that lower amounts of electroactive h (b, d :d I I I I I I~ I 7 T -T 1-1 0.7 0.5 0.3 0.1 017 0.5 0.3 0.1 E/V vs. SCE Fig. 2 Continuous cyclic voltammograms recorded at Nafion-coated electrodes immersed in 0.05 mol dm-3 H2S04 with (a) TH+, (6) MB +. Cyclic voltammograms recorded at Nafion-coated elec- trodes immersed in 0.5 rnol dmd3 H2S04 with (c) TH+, (d) MB+.Thickness of the Nafion film is 2.5 pm. 0.7 V 150 PA 0.7 v 100 mV EfV vs. SCE Fig. 3 Cyclic voltammograms of MB+ incorporated into (a) Na+-Nafion film and (b) H+-Nafion film in 0.05 rnol dm-3 H,SO,; (c) Na+-Nafion film and (d) H+-Nafion film in 0.5 rnol dm-3 H,SO,. Thickness of the Nafion film is 2.5 pm. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 1 Electrochemical data for TH' and MB' dyes incorporated into Nafion-film-coated electrodes in 0.05 rnol dm-3 H2S04 (thickness of Nafion film, 2.5 pm) electrode system 1.2 3 9 T(E2)/109rnol cm-2 DaPp(E2)/1O8cm2 s-' E~JVus. SCE T(E1)/108rnol cm-' E:,/V us. SCE D,pp(E1)/lOBcm2 s-l 4.15 6.32 0.36 1.14 0.18 1.13 6.22 9.82 0.36 0.83 0.19 3.43 1.30 5.63 0.35 1.86 0.13 1.12 5.18 8.72 0.38 0.83 0.14 2.25 I, Pt/H+-Nf/TH+;5 Pt/Na+-Nf/TH+;3, Pt/H+-Nf/MB+;& Pt/Na+-Nf/MB+.El, usual redox wave; E2, unusual redox wave. species are adsorbed by Na+-Nafion than by H+-Nafi~n;~' this is reflected in Table 2. The diffusion coefficients (D,,,), or charge transport, of TH+ and MB+ incorporated into the H+-Nafion and Na+- Nafion films were measured in 0.05 and 0.5 mol dm-3 H2S04 at different dye concentrations in the film, and the data are given in Tables 1 and 2, respectively. To determine the diffusion coefficients for dyes at Nafion films by chrono- amperometric experiments, the potential was stepped from +0.7 to +0.3 V and from +0.3 to -0.1 V us.SCE in 0.05 mol dm-3 H,S04, and for 0.5 mol dm-3 H2S04 the poten- tial was stepped from +0.7 to -0.1 V us. SCE. The diffusion coefficients for the dyes in the Na+-Nafion film are much higher than those for the dyes in the H+-Nafion film. This is because the concentration of the electroactive species in the Na+-Nafion film is lower than that in the H+-Nafion film. It has already been reported" that the diffusion coefficients decrease with increasing concentration of dyes in the film. The total amount of electroactive species determined for the two redox reactions in 0.05 mol dm-3 H2S04 is almost equal to the amount of electroactive dye molecules measured in 0.5 mol dm-3 H2S04.This observation shows that the electro- active dye molecules adsorbed by the interfacial region are transported immediately to the ionic cluster region at higher acid concentrations.The concentration of MB+ adsorbed by the Nafion film was varied from 1.30 x lo-' mol cm-2 to 1.14 x lo-* mol cm-2 and cyclic voltammograms in 0.05 mol dm-3 H2S04 were recorded and are shown in Fig. 4(a) and (b).When the amount of MB+ adsorbed in the Nafion film is low (1.30 x lo-' mol crn-'), a less intense unusual redox wave appeared. At higher concentrations of MB+ (1.14 x lo-* mol cm-2), the unusual redox wave was more intense. These observations show that at lower concentrations of MB', the dye molecules are mostly adsorbed by the sulfonate ionic cluster regions. At higher concentrations of MB +,the dye molecules are adsorbed by the interfacial region as well as the ionic cluster region of the swollen Nafion film.The swollen Nafion film consists of solvated -SO, head groups and spherical counter-ion-solvent clusters of 40 A diameter, resembling a reverse micellar structure with a hydrophobic fluorocarbon chain.".' 1*30*31 Since the inter- facial region contains higher concentrations of fluorocarbon chain, the interphase is more hydrophobic and the interfacial Table 2 Electrochemical data for TH' and MB' dyes incorporated into Nafion-film-coated electrodes in 0.05 rnol dmP3 H2S04 (thickness of Nafion film, 2.5 pm) electrode system' r/108 mol cm-2 Ep,/V us. SCE DaPd1O8cm2 s-l 1. 2 1.25 0.98 0.25 0.24 1.35 3.84 3 1.89 0.19 1.98 4 1.04 0.19 3.80 a As for Table 1.A .o LI 0.7 E/V vs. SCE Fig. 4 Cyclic voltammograms of MB+ incorporated into a Nafion film of thickness 2.5 pm in 0.05 rnol dm-3 H2S04: (a) r = 1.30 x mol cm-2 and (b)r = 1.14 x rnol cm-' region has a lower void volume than the cluster region.30 In the present study, in 0.05 rnol dm-3 H,SO, the dye mol- ecules enter into the film and preferentially reside in the higher void volume of the cluster region to give the usual redox wave. At increased concentrations, the dye molecules also enter into the lower void volume of the interfacial region to give the unusual redox wave. The ionic cluster region is well solvated at higher acid concentrations.Under these conditions the dye molecules reside only in the well solvated ionic cluster region to give the usual redox wave. This study clearly establishes the influence of acid concentration on the nature of the Nafion film and consequently on the electro- chemical behaviour of the incorporated low-charged dye molecules. Financial support from the Department of Science and Tech- nology and the Department of Atomic Energy, Government of India is gratefully acknowledged. References 1 T. D. Gierke and W. Y. Hsu, in Perfluorinated Ionomer Mem- branes, ed. A. Heisenberg and H. L. Yeager, Am. Chem. SOC. Symp. Ser., 180, American Chemical Society, Washington DC, 1982. 2 C. R. Martin, I. Rubinstein and A. J. Bard, J. Am. Chem. SOC., 1982,104,4817.3 Y. M. Tsou and F. C. Anson, J. Phys. Chem., 1985,89,3818. 4 T. P. Henning and A. J. Bard, J. Electrochem. SOC., 1983, 130, 613. 5 N. E. Prieto and C. R. Martin, J. Electrochem. Soc., 1984, 131, 751. 6 D. A. Buttry and F. C. Anson, J. Am. Chem. SOC., 1983,105,685. 7 J. Leddy and A. J. Bard, J. Electroanal. Chem., 1985,189,203. 8 D. R. Lawson, L. D. Whiteley, C. R. Martin, M. N. Szentirmay and J. I. Song, J. Electrochem. SOC., 1988, 135,2247. 9 W. Grot, Chem. Id.,1985,647. 10 H. L. Yeager and A. Steck, J. Electrochem. SOC.,1981,128,1880. 1244 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 11 12 13 14 15 G. Pourcelly, A. Oikonomou, C. Gavach and H. D. Hurwitz, J. Electroanal. Chem., 1990,287,43. K. A. Striebel, C. G.Scherer and 0.Haas, J. Electroanal. Chem., 1991,304,289. R. J. Lin, T. Onikubo, K. Nagai and M. Kaneko, J. Electroanal. Chem., 1993,348, 189. I. Rubinstein, J. Electroanal. Chem., 1985, 188,227. W. J. Vining and T. J. Meyer, J. Electroanal. Chem., 1987, 237, 191. 21 22 23 24 25 26 27 R. Memming, Prog. Surf:Sci., 1984,17,7. P. V. Kamat, J. Electroanal. Chem., 1984,163,389. P. V. Kamat and N. N. Lichtin, J. Phys. Chem., 1981,85,814. N. Oyama and F. C. Anson, J. Electrochem. SOC.,1980,127,640. H. S. White, J. Leddy and A. J. Bard, J. Am. Chem. SOC., 1982, 104,4811. K. R. Gopidas and P. V. Kamat, J. Phys. Chem., 1990,94,4723. H. Mohan and R. M. Iyer, J. Chem. SOC., Faraday Trans., 1992, 88,41. 16 17 18 19 20 R. Harth, U. Mor, D. Ozer and A. Bettelheim, J. Electrochem. SOC.,1989,136,3863. Z. Lu and S. Dong, J. Chem. SOC., Faraday Trans. I, 1988, 84, 2979. S. Kuwabata, J. Nakamura and H. Yoneyama, J. Electroanal. Chem., 1989,261,363. A. R. Guadalupe, K. E. Liu and H. D. Abruna, Electrochim. Acta, 1991,36, 881. S. A. John, K. V. Gobi, A. Ramasubbu and R. Ramaraj, Res. Chem. Zntermed., 1992,18,203. 28 29 30 31 P. D. Wildes and N. N. Lichtin, J. Am. Chem. SOC., 1978, 100, 6568. M. N. Szentirmay and C. R. Martin, Anal. Chem., 1984, J6, 1898. L. D. Whiteley and C. R. Martin, J. Phys. Chem., 1989,93,4650. R. Naegeli, J. Redepenning and F. C. Anson, J. Phys. Chem., 1986,90,6227. Paper 3/06077B; Received 1lth October, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949001241

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(9), 1245-1250 Integral Equation Theory for Associating Liquids :Highly Asymmetric Electrolytes Jun Wang and A. D. J. Haymet" School of Chemistry, University of Sydney, NSW 2006,Australia The associating liquids formalism of Wertheim is used to study association in highly asymmetric model electro- lytes, with charge asymmetry + 1 : -10, +I : -20 and size asymmetry 1 : 13 and 1 :8. These models are designed to capture some features of ionic colloids, micelles and proteins, but fall short of a complete descrip- tion. Explicit predictions are made for the fraction of associated species in the electrolyte. The theory yields improved agreement with structural data from existing computer simulations. Moreover, it predicts correctly that a large fraction of the large, highly charged ions are associated. However, the current level of theory is not a complete description of association.1. Asymmetric Electrolytes This paper studies simple models of highly asymmetric elec- trolytes using the associating liquids formalism of Wertheim' adapted by us to treat molecules interacting uia continuous, spherically symmetric potential^.^.^ Previously we have studied exactly the same models using the hypernetted chain (HNC) integral equation4 and bridge function corrections to HNC (HNC + B).' The motivation for this study is stated simply. Solutions of polyelectrolytes such as proteins, micelles and ionic colloids may behave quite differently from simple electrolytes, due in part to the asymmetry in charge between the polyelectrolyte molecules and the surrounding counter-ions and co-ions in the solution.Our view4 is that the dynamical properties of the system, including reaction rates, can be understood only if the spatial correlations between all ions in the solution are known. There are many alternative approaches to this problem6 and we do not intend to review these here. Instead, we address just one feature of these systems, which has attracted little interest to date, namely the association of ions in solu- tion. A simple model is all that is needed to see this feature. We consider a highly charged spherical electrolyte, denoted for example XZ0-, surrounded by monovalent counterions A+. (In this calculation the concentration of co-ions B-is zero.) Our goal is to predict the equilibrium concentration of all possible associated species, such as the 'dimer' X2'-A+, the 'trimer' X20-(A+)2, and higher-order 'n-mers' X2'-(A'),-'.Most researchers seem to feel intuitively that in solution the polyions exist not as 'bare' ions but rather as highly associated groups of ions which are more nearly elec- trically neutral. We agree with this proposal and this paper reports our partial progress in addressing this goal. A weak point in our treatment is the approximation of the aqueous solvent by a structureless dielectric continuum. Else- where we report our treatment of atom-based models of water by molecular dynamics (MD) simulation' and integral equation theories.* Nevertheless, since our goal is to study ion association, and the present simplified model of the poly- electrolyte is capable of studying this effect, we present our results and leave for the future the unification of this approach with a molecular-level description of the solvent.Our results are presented in Section 4 and compared with those from Monte Carlo (MC) simulations for + 1 : -10 electrolyte by Vlachy et aL4 and with those from a MD-reference HNC (RHNC) procedure for + 1 : -20 electrolyte by Lin~e,~q" in which MD simulations for the reference system and the RHNC perturbation method have been com- bined. 2. Model Asymmetric Electrolyte The interaction between ions is given by4." where the subscripts a and b denote + for counter-ions with z, = 1 and -for polyions with valences z-= -10 and -20, gab = O, + Ob, b-' = kT, k is the Boltzmann constant and T the absolute temperature. The parameters A,, and go used in this work are those used by Vlachy et aL4 and Linse,' and are listed in Table 1.The solvent is approximated by a dielectric continuum with a Bjerrum length Be'/& x 7.15 A, to model4 aqueous solutions at T = 298 K.This model mimics salt-free aqueous solutions : the free amphiphiles present in real micellar solutions are neglected. The potential energies are plotted in Fig. 1. We note that there are two length scales in each of the potentials. The asymmetry in size is ca. 1 : 13 for the + 1 : -10 electrolyte and 1 : 8 for the + 1 : -20 electrolyte.3. Brief Summary of Associating Electrolytes Associating Liquids Theory Ion clustering in highly asymmetric model electrolytes is treated by the following the~ry.~.~ Ion association follows from the separation of the pair potential, where U$)(r) and U$!(r) are the non-associating and associ- ating parts of the potential,' and Go, is the Kronecker delta Table 1 Potential parameters for + 1 : -10 and + 1 : -20 electro-lytes 1 : -10 0.08299 0.7 9.15 18.04 18.04 18.04 1 : -20 0.09960 2.0 15.0 0.795 15.89 371.876 1246 100 90 80 70 60 I-% 50'40 30 20 10 0 &=--_________-_-------___---I I I 1 I-10 0 20 40 60 80 100 r/A Fig. 1 Potential energies of the model electrolytes studied here: +1 :-10 (-) and +1 : -20 (---) electrolytes function.Unlike the strongly associating liquids, where dimers and trimers are formed by 'chemical bond^',^*^^*^^ eqn. (2) introduces a mechanism for the association of unlike ions. Using the 'energy definition',2 the associating part of the potential energy between unlike ions U?)-(r) is written as where U, is determined by minimizing the Helmholtz free energye2v3 OrnsteieZernike Equation The total correlation functions, h$(r) =g$(r) -hOiSoj, are determined from the analogue of the Ornstein-Zernike (OZ) equation, 'ab = cab +1cad * Pd * 'db (4) where * denotes convolution. The matrices are defined to be (5) where the subscript 0 indicates a non-associated species, and the subscript 1 an associated species. The number density for species a is pa = zp:-1, where pi-(for n = 1, 2, 3, ...) denotes monomer, dimer, trimer and higher 'n-mer 'densities.Eqn. (4) is renormalized as usual to treat the long-range Coulomb intera~tion.~.~ This renormalized OZ equation in Fourier space may be written W)= {I-cz + &(k)Plm)PI -{CZ+&(k)Pl6Qk) x 1+ +P6Wl -Se(W (6) where Z is the 4 x 4 unit matrix. The matrix M(k) = 6@k) -Sqk) is obtained from the Fourier transforms of 6H(r) and SC(r).The explicit forms of matrices p, SH(r) and 6C(r) are given in our study of the +1 : -3 ele~trolyte.~ J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 The short-ranged functions 6h$(r) and 6c$'(r), the elements of the matrices 6H(r)and 6C(r),respectively, are defined to be 6c$(r) = c$(r) +dOi6, flu:& (7) and ah;;@) = h$(r) -q$(r) (8) where U:&) is the Coulomb part of the total potential, uab(r), and q;$r) = 6oi 6oj q"DbH(r) where qgH(r)is the Debye-Huckel (DH) potential of mean f~rce.~,~ Approximate Closure We now introduce the approximation into the theory. First, we assume that the ions can form only dimers and trimers and examine the consequence of this approximation.Note that there are two distinct possible trimers in the system. We use the symbol +-+ to denote the trimer consisting of one polyion and two counterions and -+-for the other case. Secondly, we adopt the renormalized HNC-like closure rela- tions used earlier,2.3 6c$\(r) = g$)(r) -q"db(r)-Gz$\(r) -1 Sc;;b,(r) = Gz;;b,(r)[g"obo(r) -11 dcll,\(r) = 6z",(r)[g"db(r) -13 6cyb1((r) = [6z';b,(r)6z",b,(r) +Z"lbll(r) +(1 -aab)f$)-(r) + Kb(r)lg"obo(r) -6z?l(r) (9) where s"db(r)= expC-BU!$(r) + 6z"dbW +q"db(r)I fy)-(r) = exp[ -fiUy)-(r)] -1 and U!j(r) = Ub")(r)-Uf!(r) The functions account for trimer formation.For asym- metric binary mixtures, under the superposition approx-imation for trimers, we find qb(r) = (l -6ab)f$'-(r)[pt +p: I:(r)]/r + (l -6ac)p', 12(r)/r (10) C and the monomer densities pof and pi are determined self- consistently from the equations p+ = pof +4~pofpiJ+4~(pof)~pkJ++2npL(p,)'J-p-= pi +47~pofpiJ+47~(pi)~p;J-+27r~i(p;)~J+ (11) The division energy, U,, in eqn. (3) is determined by mini- mizing the Helmholtz free energy, given by fl(A -A("))/V= C [pa ln p:/pa +(pa -p",/2 a +nn(p"02 1(l -6ab)Pb,Ja1 (12)b where A'") is the Helmholtz free energy of the reference system with the interaction U$)(I-).~The fractions of dimers and trimers are obtained from the following expressions x2 = 8VofPOJ/Pt 9 x3(+ -+)= 64PL)2POJ+/pt and x3(-+-) = 64PO)2PLJ-/Pt J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 where pt = p+ + p-. The explicit expressions for I:@), I&), J" and J in eqn. (10)-(13) are given by Wang and Ha~met.~ Eqn. (6) and (9) are solved numerically with 2048 grid points and a mesh 6r = 0.235 A by Picard iteration. The monomer densities, p;, are obtained iteratively from eqn. (11). The details of the algorithm are given by Kalyuzhnyi et aL2 and Ichiye and Haymet.14 The complete pair correlation functions are given by We note here that the monomer fractions for species a, x: = p",p" describe the proportion of 'bare' ions that exist in solu- tion.The input to the theory is the potential energy given by eqn. (1) and the total number density pa for species a. The theory predicts the structural and thermodynamic properties and the fraction of monomers, dimers and trimers. 4. Predictions for Asymmetric Electrolyte Thermodynamics, Dimer and Trimer Concentrations The reduced excess energy, E = EJp, kT, and the reduced osmotic pressure, (6, = P/ptkT, calculated via the virial route and the reduced optimal division energy, U,/kT, are collected in Table 2.The thermodynamic data, together with the HNC results, are compared with the MC results for + 1 : -10 elec- trolyte by Vlachy et aL4 and the RHNC results for +1 : -20 electrolyte by Linse.' The agreement between the present pre- dictions and the MC, RHNC results are good, although not all the thermodynamic quantities are improved over the results of HNC. A similar fact has been observed by Duh and Haymet" for a 2 : 2 electrolyte study by including 'bridge functions' into the cluster expansions for the pair correlation functions. The fractions of 'n-mer', x, (n = 1, 2, 3), are displayed in Table 3. Unlike the situation for 2 : 2 and 1 :3 model electro- lyte~,~.~where the trimer fractions are always lower than that of the dimers, the fraction of + -+ trimers is comparable to the dimer fraction for +1 : -10 electrolyte and even higher than the dimer fraction for +1 : -20 electrolyte, owing to the high asymmetries in both charge and size.Moreover, the monomer fractions for species a presented in Table 3 show that only 16% of polyions remain non-associated in the + 1 : -20 electrolyte. The -+ -trimer fractions are not displayed in Table 3 since they are four or seven orders of magnitude less than the + -+ trimer fractions for + 1 : -10. and +1 : -20 electrolytes, respectively. 1247 Complete Pair Correlation Functions and Partial Structure Factors It is well known that the HNC approximation produces unphysical descriptions for the structure and critical behav- iour of model electrolytes in the strong-coupling and low- concentration regime.'6J More specifically, the HNC approximation overestimates both counter-ion-poly-ion and counter-ion+ounter-ion correlations for highly asymmetric model electr~lytes.~~~~~~'~ This, in turn, leads to the shift of the first peak of the polyion-polyion correlation function to smaller separations than in the results from sim~lations.~~'~ Apparently, these deficiencies of HNC are because of its in- ability to account for ion association.'7 The predictions for the complete pair correlation functions and the partial structure factors from the present theory, together with the HNC results, for the + 1 : -10 and + 1: -20 model electrolytes are now stated.Fig.2-4 display the results for the + 1 : -10 electrolyte. Our associating inte- gral equation theory yields an improved counter-ion-counter-ion correlation function g+ +(r)shown by the solid line in Fig. 2. Fig. 3 shows that the peak of the counter-ion- 1.o 9++ 0.5 0.0 r/A Fig. 2 Complete pair correlation function for pairs of counter-ions for the + 1 : -10 electrolyte at c, = 0.08299 mol dm-'. (-) Present theory for the optimized division energy, U,/kT = -4.85, (---) HNC approximation; (0)MC simulation result by Vlachy et a1.4 Table 2 Thermodynamic quantities d = A,,/NkT, E = E,,/p, kT and & = P/p,kT for + 1 : -10 and + 1 : -20 model electrolytes 1: -10 0.08299 1: -20 0.0996 a Not available. Table 3 Monomer fractions x: + 1 : -20 model electrolytes 1: -10 0.082 99 1: -20 0.009 96 this work -4.85 -1.703 -2.126 0.625 aMC --.2.14 f0.05 0.625 f 0.02 HNC -1.699 -2.143 0.635 this work -6.97 -2.709 -3.152 0.596 RHNC -a -3.152 0.609 HNC -2.708 -3.161 0.606 = &/pa, for species a and the fractions of 'n-mers', x, E (pT-+ p,--l)/pt (n = 1, 2, 3), for + 1 : -10 and -4.85 0.907 0.327 0.854 0.075 0.071 -6.97 0.935 0.162 0.919 0.039 0.042 J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 8 6 g+-4 2 0 r/A Fig. 3 Complete pair correlation function for pairs of unlike ions for the +1 : -10 electrolyte at c+ = 0.08299 mol dmV3. Lines have the same meaning as in Fig. 2. polyion correlation function, g + -(r), which is overestimated by HNC, is decreased in the present work, although a little severely as is the case for the +1 : -15 model electrolyte examined using the HNC + B approximation.' The polyion-polyion correlation function, g --(r), from the present theory, exhibited in Fig.4, is very close to the result of the HNC approximation. Bearing in mind that the unlike ions can at most form trimers in the present framework, we now turn to the +1 : -20 case. Fig. 5-7 present the complete pair correlation functions for this strongly coupled electrolyte. Fig. 5 shows that the counter-ion-counter-ion pair correlation function, g+ +(r), is improved slightly over the HNC approximation result. The primary maximum in the counter-ion-polyion pair correlation function g+ -(r), shown in Fig.6, is still underestimated as for the + 1 : -10 electrolyte. The polyion-polyion pair correlation function, g --(r), is again essentially overlapping with the HNC result, as displayed in Fig. 7. The partial structure factors, So&), for the two electrolytes under consideration are given in Fig. 8 and 9, with the HNC 1.5 1.o €I--0.5 0.0 rlA Fig. 4 Complete pair correlation function for pairs of polyions for the + 1 : -10 electrolyte at c, = 0.08299 mol dmW3. Lines have the same meaning as in Fig. 2. 1.5 1.o 9++ 0.5 0.0 __ 10 20 30 40 rlA Fig. 5 Complete pair correlation function for pairs of counter-ions for the + 1 : -20 electrolyte at c, = 0.0996 mol dm-3. (-) Present theory for the optimized division energy, U,/kT = -6.97; (---) HNC result; (0)result from an RHNC procedure by Linse." results also included.The results reported above suggest that ion association is still not being fully addressed within the present theory, in which unlike ions can only be associated into dimers and trimers. Further discussions to elucidate this point are given below. Partial Pair Correlation Functions One important concept in Wertheim's two-density integral equation theory is 'monomer depletion',' that is the process whereby the strong attraction between particles leads to association thus decreasing monomer densities. Mathemati-cally, the difficulty encountered by the HNC approximation for strongly coupled systems is due to the deep well in the attractive potential, or equivalently the large value of the Mayer function.Unlike the other routes used to unravel this problem before, such as including the 'bridge functions' in the HNC + B approximati~n~~'~ and introducing a three- body potential in HNC+3 theory," the depletion of mono- r/A Fig. 6 Complete pair correlation function for pairs of unlike ions for the + 1 : -20 electrolyte at c, = 0.0996 mol dm-'. Lines have the same meaning as in Fig. 5. J. CHEM.SOC. FARADAY TRANS., 1994, VOL. 90 0.0I,,,,,,,,,,,,B I I I 10 20 40 60 80 100 120 140 r/a Fig. 7 Complete pair correlation function for pairs of polyions for the + 1 : -20 electrolyte at c, = 0.0996 mol dm-'. Lines have the same meaning as in Fig. 5. I I I I I I , I 0.0 0.1 0.2 0.3 0.4 0.5 q/A-' Fig. 8 Partial structure factors for the +1 : -10 electrolyte at c+ = 0.08299 mol dm-'.(-) Present theory with the optimized division energy, U,/kT = -4.85; (---) HNC results. I 1 1 I I I I I-0.5I I 0.0 0.1 0.2 0.3 0.4 0.5 9lA-l Fig. 9 Partial structure factors for the + 1 : -20 electrolyte at c, = 0.0996 mol dm-'. (-) Present theory with the optimized division energy, U,/kT = -6.97; (---) HNC results. 1249 mers serves to suppress the catastrophe due to a large value of the Mayer function in Wertheim's two-density integral equation formalism.' To illustrate the importance of monomer depletion, we present the partial pair correlation functions, g$'(r), and the corresponding complete pair corre- lation functions, gob(r),for the + 1 : -10 electrolyte in Fig.10-12. Fig. 10 shows that the counter-ion-counter-ion correlation is mainly determined by the correlation between non-associated counter-ions g&+(r). Monomer depletion results in a decrease of g&,+(r) and the consequent improvement of the complete correlation function, g + +(r),over the HNC result. The partial correlation function, g&+(r), would be further reduced if more monomers are depleted. The results in Fig. 11 imply that the principal peak in the complete counter-ion-polyion pair correlation function, g+ -(r), arises predominantly from the contribution of the partial pair correlation between associated unlike ions, g:l-(r), which will be enhanced as more ion clusters are 0 10 20 30 40 50 rlA Fig. 10 Partial pair correlation functions for the + 1 : -10 electro-lyte: (-4 g++(r),(----) 90+0+('),(---) 90+l+(r),(-*-I 9:m 10 85 0 0 5 10 15 20 25 r/A Fig.11 As Fig. 10, but for the counter-ion-polyion correlation: (. . * *) 9;0-(') 0 20 40 60 80 100 120 r/A Fig. 12 As Fig. 10, but for polyions formed, i.e. more monomers are depleted. Fig. 12 shows the case for the complete and partial polyion-polyion pair corre- lation functions, where g&,-(r) plays the main role in deter- mining g --(r). 5. Conclusion Ion association in + 1 : -10 and + 1 : -20 model electro- lytes is investigated by Wertheim’s associating liquids integral equation formalism. It is predicted that only a small portion of ‘bare’ polyions exist.The effect of monomer depletion on the correlations in the + 1 : -10 electrolyte is analysed by examining the partial pair correlation functions. However, comparisons with the ‘exact’ results reveal that ion associ- ation is still not addressed sufficiently in the present formal- ism, in which ions can form at the most trimers. We believe that it is necessary to consider at a minimum the formation J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 of tetramers, consisting of one polyion and three counterions, to improve the present formalism. Graphically, this means that four-particle diagrams will appear in the density equa- tions and closure relations. This research was supported by the Australian Research Council (ARC) (grant No.A29131271). A.D.J.H. acknow- ledges gratefully many helpful conversations on this topic with Prof. V. Vlachy, Dr. Yu.Kalyuzhnyi, Prof. M. Holovko and Mr. M. J. Booth. References 1 M. S. Wertheim, J. Stat. Phys., 1984,35, 19, 35. 2 Y. V. Kalyuzhnyi, M. F. Holovko and A. D. J. Haymet, J. Chem. Phys., 1991,95,9151. 3 J. Wang and A. Haymet, J. Chem. Phys., 1994,100,3767. 4 V. Vlachy, C. H. Marshall and A. D. J. Haymet, J. Am. Chem. Sac., 1989,111,4160. 5 J. ReSEiE, V. Vlachy and A. D. J. Haymet, J. Am. Chem. SOC., 1990,112,3398. 6 Y. N. Vorobjev, J. A. Grant and H. A. Scheraga, J. Am. Chem. SOC.,1992, 114, 3189. 7 A. Nyberg, D. E. Smith, L. Zhang and A. D. J. Haymet, J. Chem. Phys., 1994, submitted. 8 D-M. Duh, D. Perera and A. D. J. Haymet, J. Chem. Phys., 1994, in preparation. 9 P. Linse, J. Chem. Phys., 1990,93, 1376. 10 P. Linse, J. Chem. Phys., 1991,94,3817. 11 P. J. Rossky, J. B. Dudowicz, B. L. Tembe and H. L. Friedman, J. Chem. Phys., 1980,73,3372. 12 H. C. Andersen, J. Chem. Phys., 1973,59,4714. 13 H. C. Andersen, J. Chem. Phys., 1974,61,4985. 14 T. Ichiye and A. D. J. Haymet, J. Chem. Phys., 1988,89,4315. 15 D-M. Duh and A. D. J. Haymet, J. Chem. Phys., 1992,9,7716. 16 H. L. Friedman, Annu. Rev. Phys. Chem., 1981,32,179. 17 J. S. Hsye, E. Lomba and G. Stell, Mol. Phys., 1992,75, 1217. 18 B. B. Laird, J. Wang and A. D. J. Haymet, Phys. Rev. E, 1993, 47.2491. Paper 3/07172C; Received 6th December, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949001245

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(9), 1251-1259 1251 Evanescent Wave Spectroscopy :Application to the Study of the Spatial Distribution of Charged Groups on an Adsorbed Polyelectrolyte at the Silica/Water Interface Mathias Trau,f Franz Grieser, Thomas W. Healy and Lee R. White* School of Chemistry and Department of Mathematics, University of Melbourne Parkville Australia 3052 A new evanescent wave experimental technique with the capacity to determine simultaneously the spatial dis- tribution of several chromophoric species located near a reflecting interface is reported. In a first step, to illustrate the capacity of this technique, the absorption of a model polyelectrolyte polymer (polyacrylamide/ diacetone acrylamide copolymer grafted with an ionizable acridine chromophore) (EPI-26) onto the silica/ aqueous solution interface has been studied.For this sytem, it has been demonstrated that variable angle of incidence evanescent wave spectroscopy may be used to determine quantitatively both the surface excess, r, and the mean separation distance from the interface 2, of charged and uncharged segments attached to the adsorbed polyelectrolyte. The technique has also been used to measure the kinetics of adsorption, as well as changes in r,2 and degree of ionization of the adsorbed layer as a function of surrounding solution conditions (e.g. pH or salt concentration). It was found that increasing the ionic strength resulted in a larger surface excess of the polyelectrolyte, and decreasing the pH, which further ionizes the polyelectrolyte, reduced the surface excess and caused the polyelectrolyte layer to expand.Both of these trends are in accordance with what is expected from simple electrostatic considerations and further show the sensitivity of the evanescent wave tech- nique to the microscopic structure of such adsorbed layers. The determination of the conformation and charge distribu- 28-3 1. Although ellipsometry, in principle, can give confor- tion of adsorbed polyelectrolytes is an area which has great mational information about the adsorbed layer,32 it is gener- importance in both colloid science'*2 and molecular ally limited by sensitivity constraints to measuring only the Many review articles outlining the current experi- surface excess of the adsorbed polymer,*' with the exception mental and theoretical approaches to this problem are avail- of highly reflective substrates (e.g.metals) where an able in the literat~re.~-'' Although, at present, there exists a 'ellipsometric thickness', a measure of the extension of the relative abundance of theoretical work in this area, in partic- polymer normal to the interface, can also be measured." ular calculations uia the lattice mean-field approach, e.g. ref. None of the existing experimental techniques is currently able 11-14, as well as various Monte Carlo calculations, ref. 15 to provide the conformational, charge distribution and and 16, there still remains a great paucity of experimental kinetic information about the adsorbed polyelectrolyte layer data on well defined systems which can critically test these that is needed for a full understanding of the polyelectrolyte theoretical predictions.The experimental determination of adsorption process on a molecular level. The ideal technique adsorbed polymer conformation has been previously should be able to produce quantitative, real-time, in situ data attempted uia a wide variety of techniques: e.g. neutron about each of these microscopic features of the adsorbed scattering/reflectivity,' 9 ' ellipsometry, 9 2o photon corre-layer. All of the above experimental techniques are an lation spectroscopy,' NMR,2' EPR22-25 and capillary flow approximation to this optimum. techniques.26 Of these, only the neutron scattering/reflectivity In view of these limitations, we have developed a new spec- techniques provide a method of rigorously determining the troscopic technique, variable angle of incidence evanescent conformation of the entire adsorbed polymer layer (usually wave spectroscopy (VIEWS), which has the capacity to known as the polymer segment density profile); however, measure the spatial distribution of several spectroscopically they give no information about the charge distribution in the distinct chromophores near a reflecting interface simulta- adsorbed layer.The neutron techniques also suffer from a neously. VIEWS measures the attenuation of a totally inter- number of practical disadvantages which has limited their nally reflected beam of light which results from the use: e.g.(i) the necessity of a strong neutron beam; (ii) the interaction of the evanescent wave, generated at the solid/ need to prepare deuteriated compounds synthetically; (iii) the aqueous solution interface, with absorbing chromophores difliculty of obtaining kinetic data because of the weak scat- attached to the backbone of the adsorbed polyelectrolyte. tering; and (iv) difficulties with the inversion of raw Conformational information can be obtained by varying the scattering/reflectivity data into unique segment density pro- depth of penetration of the evanescent wave, primarily con- file~.'~ trolled by reflection angle, and charge distribution informa- For these reasons, the majority of experimental work tion can be obtained by varying the wavelength of the beam which has been previously reported in this area has mainly and probing ionizable chromophores which possess clearly been concerned with techniques which measure only the total distinct spectra in the ionized and in the neutral forms.adsorbed amount (surface excess) of the adsorbed poly-Although we shall report results here for a model poly- electrolyte layer and provide no information about the electrolyte, one which has been specifically synthesized to microscopic structure of the adsorbed polyelectrolyte, e.g. ref. contain an ionizable probe (acridine) randomly distributed along the backbone, the techniques is not restricted to such artificially tagged polymers. In principle, any chromophore t Present address : Department of Chemical Engineering, Prin- ceton Materials Institute, Princeton University, Princeton, New (or chromophores) on the polymer backbone may be used, as Jersey 08544, USA.long as the distribution of the chromophore on the backbone 1252 is either known or random. As well as being able to deter- mine simultaneously the spatial distribution of more than one chromophore in the interfacial region, our technique has other significant advantages over evanescent wave fluores- cence techniques which have been previously reported, e.g. ref. 33. These include: (i) the problem of fluorophore quantum yield variation with distance from the is completely circumvented ; (ii) fluorescence induced by surface and/or background scattered light (an annoying arti- fact which is extremely difficult to remove accurately from the data) is completely absent in this technique; and (iii) the tech- nique does not necessarily require a fluorescent (or absorbing) probe to be attached artificially to the polymer backbone.In this publication, emphasis is placed on correcting all VIEW spectra for scattering effects which are observed for certain types of adsorbed polymer layers. A novel multiple- wavelength analysis scheme, which may be used to extract concentration profile information for more than one (spectroscopically distinct) chromophore in the adsorbed layer, is also presented. Experimental Apparatus The apparatus is shown schematically in Fig. 1 and 2. All of the components displayed in these diagrams fit inside the cavity of a standard UV-VIS spectrophotometer (Cary 2215, Varian).The incoming sample beam from the spectropho- tometer is passed through a collimating lens (designed to col- limate the beam to a maximum divergence of 0.25"), a Glan-Thompson polarizer (which can be set for either s or p polarization states), a collimating slit and a silica waveguide (where, depending on the chosen angle of incidence, the beam is totally internally reflected from 5 to 20 times) before being passed through to the spectrophotometer detector viu a series of mirrors. The angle of incidence, for all of the total internal reflections inside the waveguide, is accurately controlled via a precision rotor which can rotate the waveguide and both mirrors around a central pivot point to within a precision of 1 arc min [see Fig.2(u)]. This geometry allows accurate control of the angle of incidence while preserving the natural path of the beam to the spectrophotometer detector at all angle settings. The incoming reference beam is passed through a Glan-Thompson polarizer, identical to the one in the sample beam, and attenuated by a neutral density filter. Polarizer neutral density filter mc;;ingisample sl~ r' detector+ ,I ,tobeam mirror 2 polarizer pivot point Fig. 1 Schematic illustration of the variable-reflection-angle ATR apparatus. All of the optical components displayed fit inside the sample cavity of a standard laboratory spectrophotometer (Cary 2215, Varian).The waveguide and the two mirrors are mounted on a precision rotary table and may be rotated around the central pivot point (shown) to set precise the angle of incidence, 8,for all of the reflections inside the waveguide. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 aqueous solution4 stainless steel mounting table point c3precision rotary stage (b) aqueous solution waveguide ueous ionlsilica O-ring contact area aqueous solution I entry O-ring seal (Teflon coated) Fig. 2 (a)Assembly of the ATR cell holder and mirrors on the pre- cision rotary stage. (b) Assembly of the ATR cell holder. (c) Aqueous solution/waveguide contact area. Fig. 2(b) illustrates the flow-through mechanism by which an aqueous solution from a remote reservoir is brought into contact with the waveguide and Fig.2(c) shows the area of contact between the aqueous solution and the silica wave- guide reflecting surface. With this arrangement absorbance attenuated total reflectance (ATR) spectra may be collected for any angle of incidence in exactly the same manner as standard transmission spectra. Polyelectrol yte The polyelectrolyte used in these experiments (EPI-26) was specifically synthesized to contain an ionizable spectroscopic probe (acridine) randomly distributed along a water-soluble polymer chain (polyacrylamide/diacetone acrylamide copoly- mer) (80 mol% acrylamide, 20 mol% diacetone acrylamide). The polyelectrolyte backbone has the following schematic formula : C=0 C=0 C-0 I I NH NH I I CH3-C-CH3 CH3-C-CH3 I I p2 I c=o I CH3 The synthesis procedure for this polyelectrolyte has been described in a previous p~blication.~' The acridine probe (P) was specifically chosen to be the only ionizing group on the J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 14.01 , A 260 340 420 500 wavelength/nm I B 1.o 8.4 5.6 2.8 I 1OJ , I , , , ,235 245 255 265 wavelength/nrn Fig. 3 Absorption spectra of 9-(hydrazinoformy1)acridine in water collected at pH values of (a) 1.90,(b) 2.54, (c) 3.32, (6)3.53, (e) 3.72, cf) 4.04, (9) 4.57, (h) 5.14, (i) 5.53 and (j)6.32. A, Entire UV-VIS spec-trum. B, Expanded view of the spectral region where I, P and 11, PH+ are most clearly distinct.The isosbestic point may be clearly seen at 1= 253 nm. backbone because its spectrum is clearly distinct from that of its conjugate acid, the acridinium ion (PH') (see Fig. 3). The average molecular weight of the polyelectrolyte was esti-mated from the synthesis conditions40 to be ca. 1oOOOO and UV spectroscopy was used to determine the total amount of grafted acridine on the polymer backbone (12.3 monomer molOh). Waveguides The waveguides used in this study were all prepared from Supracil glass (vitreous silica) supplied by H. A. Groiss Ltd. Plate dimensions were 5 cm x 2 cm x 0.2 cm with the short edges cut precisely to an angle of 70" (see Fig. 1 and 2). The polishing and cleaning procedures used to prepare the reflec- ting surfaces have been described el~ewhere.~'.~' Theory The basic principle of this technique is to measure the attenu- ation (absorbance) of a totally internally reflected spectro- photometer beam which results from the interaction of the evanescent wave, generated at the silica/adsorbed poly-electrolyte/aqueous solution interface, with absorbing chromophores attached to a polyelectrolyte backbone. In the case where only one absorbing chromophore is present in the polyelectrolyte layer, a simple expression for the ATR absorbance, AATRA, e), can be derived from a solution of Maxwell's equations :39,41 where A is the spectrophotometer wavelength, 8 is the angle of incidence, n(0)represents the effective number of reflections the spectrophotometer beam experiences inside the silica waveguide for any chosen 8 (this parameter can be evaluated either by ray tracing or via a simple calibration technique?), &(A) is the molar absorption coefficient of the absorbing chromophore attached to the polymer backbone, p(z) is the concentration of absorbing chromophores at a distance z normal to the interface, Z,(B)/cos 8 represents the intensity of the evanescent wave at the interface (for unit incident amplitude) and 5 represents the penetration depth of the eva- nescent wave into the aqueous phase and is defined as where n, and n2 are the respective refractive indices for silica and the aqueous phase andf(8) is a function which depends only on 8 (given that n, and n2 are fixed for the experiment).A solution of Maxwell's equations for this system gives the following expressions for I,(e) :41 (3) for an s-polarized spectrophotometer beam and 4n21 cos2 e(2 sin2 8 -1122,) I,(@ = (1 -t~;~)[(1+ n$,)sin2e -n;l~ (4) for a p-polarized spectrophotometer beam (n2,= n2/n1). Provided the chromophore is distributed either randomly or with a known distribution along the polymer backbone, the segment density profile of the polymer can be inferred directly from p(z). In the case where it is randomly distrib- uted, as for the polymer reported here, p(z) corresponds exactly to the polymer segment density profile. A method for inverting eqn. (l), in order to determine p(z) from the raw data has been described in a previous p~blication.~~ Determination of Concentration Profiles for Several Chromophores Simultaneously In the case where more than one chromophore is present in the adsorbed layer (e.g.chromophores P and PH+ in poly- electrolyte EPI-26), eqn.(1) must be modified by replacing @)p(z) with E~(A)P~(Z) + e2(R)p2(z)+ . where &,(A) and p,(z),a, respectively, represent the molar absorption coefficient and concentration profile of the ith chromophore in the adsorbed layer. The data inversion technique in this case is thus com- plicated by the fact that each ATR component spectrum must be rigorously deconvoluted from the combined spectrum before the inversion process is attempted. In the following treatment, for reasons of simplicity, we shall consider only the case of two chromophores.This treatment, however, can be easily extended to any number of chromophores which possess clearly distinct spectra. The above substitution yields x exp(-z/<) dz (5) which may be rewritten in a simpler form 6) = &1(4F1(0+ &2(4F2(5) (6) where A(A, 0) = (A,& 0)cos 8/[n(8)l0(B)]}is the normal- ? The calibration technique involves recording ATR spectra, A,,&, O), for a system where p(z) is known (e.g. a bulk solution of free probe) and solving eqn. (1) for n(O). Consistent results have been obtained41 using a wide variety of water-soluble probes and these are the n(O)data which are used here. ized spectrophotometer absorbance and m =FX~) =/i(zbxd -2,~) dz (7) Because 5 changes with A, for any fixed 8, the F,(t) terms cannot be determined from a simple linear regression fitting of eqn. (6) to the A(A,8) data.To allow for this we write F~CV(~)I= Fi[V(e) + (1-X)f(e)l (8) = ~~[;if(e)]+ (A -i)f(e) (9) where 1is defined as a mid-range value of A. Substitution of this into eqn. (6) thus gives A(A 8) = &l(A)Fl(f)+ &2(4F2(0 (1-3 f aF,(f)+ q(A) 7-at A linear regression fit of A(A, 8) to a functional form of (A -1)A(A, 8) = E~(A)A+ c2(A)B+ &,(A) -A C (A -1) * -+ 4)-A D + will thus yield the parameters A = F,(f) B = F2(P) C=f-aF,(f) at D=f-aF2( f)and at for any angle of incidence 8 and for f = If(8). Note that parameters C and D will be extracted from such a fitting only if the wavelength range of the data is sufficiently large and if the level of experimental noise is extremely low.Given a finite level of experimental noise, standard least-squares method^^',^^ may be used to determine at what point eqn. (11) should be truncated. -Eqn. (7) shows that Fi(ois a Laplace transform of the con- centration profile for the ith chromophore, p&). In principle, therefore, it should be possible to determine all of the chromophore concentration profiles completely by per-forming a_" inverse Laplace transform on each Fi(f),provided each F,(t) has been measured over a wide range of evanes- cent wave penetration depths. This can, in principle, be carried out via the above apparatus by collecting ATR spectra for a variety of reflection angles [eqn.(2) shows that 5 is extremely sensitive to 81. In practice, however, the accurate inversion of Laplace transforms is an extremely dificult process which requires virtually noise-free data to distinguish between different types of profiles which possess almost iden- tical Lapla_ce transforms. An alternative approach is to rewrite F,(t) in terms of the moments for the chromophore distribution, e.g. where Ti is the surface excess of the ith chromophore and 2; J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 is the nth moment of the distribution and is defined by 1 fm = Jo dzz"p,(z) In such a representation Zi is the mean separation distance of the ith chromophore from the interface, 2: represents the mean-squared distance, 2; represents the mean-cubed dis- tance aid so on.These moments may now be extracted from the F,(<) data by performing a second linear regression fit involving eqn. (13). As with the first linear regression, given a finite level of experimental noise, a standard least-squares te~hnique~"~~may be used to determine where eqn. (13) should be truncated. Such a fitting allows rigorous extraction of the maximum number of moments possible from the experimental data: the maximum number of moments extractable is limited by the magnitude of the noise. Clearly, the more moments that can be extracted from the data, the more information we can infer about the spatial distribution of the chromophore and hence the better we can reconstruct the complete pi(z) function.If the parameters C and D in eqn. (12) can be extracted accurately over a suitable range off, a third linear regression can be performed by using the following relationship: L4 4 L4 A linear regression analysis using eqn. (16) will thus extract precisely the same moments as those in eqn. (13), with the exception of the first term. Such an analysis may provide a useful cross check of the moments determined via the linear regression involving eqn. (1 3). Evanescent Wave Scattering Fig. 4(a) shows an example of the type of raw data which is collected via this technique. Each spectrum in Fig. 4(a)rep-resents an ATR spectrum (p polarized) of the adsorbed poly- electrolyte collected at a particular angle of incidence.In this example, all of the spectra were collected 23 h after a 250 ppm EPI-26 electrolyte solution, adjusted to pH 3.2, was introduced into the ATR cell. A problem which complicates the chromophore distribution analysis described above is elastic scattering of the evanescent wave by cluster sites (small regions of relatively high refractive index) which may form in the adsorbed polyelectrolyte layer as a result of aggregation of hydrophobic moieties (e.g. diacetone acryla- mide and acridine) attached to the polymer backbone. The mechanism of formation of such hydrophobic clusters is illus- trated schematically in Fig. 5. Evanescent wave scattering from these cluster sites manifests itself in the ATR spectrum as a gradual (monotonic) increase in absorbance with decreasing wavelength and can be clearly seen in the EPI-26 ATR spectra in Fig.*a), particularly in the wavelength regions between 280-320 nm and 220-235 nm where the acri- dine chromophore is effectively non-absorbing. This type of scattering is not observed for all adsorbed polymers (e.g. surface-adsorbed polyvinyl pyridine displays no scattering ~omponent~~).If present, the scattering component of such an ATR spectrum is usually only a minor contributor to the overall spectrum and can normally be removed via standard scattering theory : Provided the scattering sites are smaller than A and possess a refractive index that does not signifi- cantly perturb the overall refractive index of the adsorbed J. CHEM.SOC. FARADAY TRANS., 1994, VOL. 90 0.8 ~ 5 0.6-a , - e-ii0.4: 0.2- 0.0- i , , , , , , 260 340 420 500 'q(b1 wavenurn ber/nrn 0.81 n 260 340 420 500 waven urnber/nrn Fig. 4 (a) Typical example of variable-angle absorbance ATR spectra (p polarized) collected for the EPI-26 polyelectrolyte adsorbed at the aqueous solution/silica interface. The above spectra were collected 23 h after a 250 ppm solution (adjusted to pH 3.0) of polyelectrolyte EPI-26 was introduced into the ATR cell. Reflection angles for these spectra area: 0 = 66.68" (top spectrum), 67.34", 68.01", 68.67", 69.34", 70.00", 70.66", 71.33", 71.99", 72.66", 73.32", 73.98", 74.64", 75.31", 75.97", 76.63", 77.28", 77.94", 78.60" and 79.25" (bottom spectrum).(b) Resultant sca tter-corrected spectra obtained by using two terms in eqn. (20),i.e. by modelling the scattering com- ponent uia RGD scattering theory. layer,t the scattering component of the spectrum, AATR(A, O),,,,, may be described uia an equation very similar to eqn. where C,,,,,a constant dependent on the optical properties of the scattering sites, represents the total amount of scattered intensity for unit incident amplitude of the spectrophotom- eter beam and n(z) represents the distribution of the scat- tering sites near the interface. The wavelength dependence for this type of scattering may be predicted from standard scattering theory44 by writing an expression for the parameter C,,,,in eqn.(17). For scattering sites smaller than the spectrophotometer wavelength, an expression for C,,,,can be derived from Rayleigh-Gans- Debye (RGD) theory1 time Fig. 5 Schematic illustration of hydrophobic cluster formation in an adsorbed polymer layer t Both of these assumptions are extremely reasonable for the aggregated clusters considered here, given that their size will be <A/10,41*46their refractive index will be c1.4 (the approximate refractive index of the bulk polymer) and, because of the small amount of scattering measured, they should be relatively dilute in the adsorbed layer. $ Ref. 44, eqn. (3.2.24). where c1 represents the polarizability of the scattering site, c0 is the permittivity of free space and P(Q) (known as the form factor) represents the modification of the scattering intensity due to the finite size or non-sphericity of the scattering sites.For scattering sites of arbitrary shape, it has the following form:45 where Q = (4nn,/A)sin(BS,,J2) is the magnitude of the scat- tering vector, f3,,,, is the scattering angle and rG is the radius of gyration of the scattering site. Higher-order terms in this equation, which start at (Qr,-J4, are usually ignored unless rG = 2 (> 200 nm for the ATR spectra considered here). Sub- stitution of eqn. (18) into (17) yields where q(0) are all constants which depend on the optical properties of the scattering sites, on n(z) and on the angle of incidence. For any ATR spectrum which exhibits evanescent wave scattering, the q(0) constants may be determined by fitting eqn.(20) to purely scattering wavelength regions of the ATR spectrum (e.g. for EPI-26, between A = 280-320 and A = 220-235 nm) where the absorbing chromophores are effectively non-absorbing and the measured scattering com- ponent is significant. An identical linear regression technique to the one suggested in the previous section can be used for this fitting. For the EPI-26 ATR spectra, a maximum of two terms in eqn. (20) is enough to obtain an adequate fit of the scattering component in all spectra. Once the q(0) values have been obtained for every angle of incidence scanned, the 'scatter-corrected' ATR spectra, AATR(A, O),o,r ,may be calcu- lated via AATR(J-9 @corr = AArR(2, 8) -AAd5 @scat (21) These scatter-corrected spectra may now be analysed via the procedure described in the previous section. If rG > A, the evanescent wave may be severely distorted by the presence of the scattering sites and result in a much more complicated theoretical situation.Provided that an adequate fit of the scattering component can be obtained via eqn. (20), this situation need not be considered. Fig. 4(b) shows the results of the above scattering correction applied to the raw ATR spectra shown in Fig. 4(a). In this example, two terms in eqn. (20) were required to obtain an adequate fit of the scat- tering component and the resulting corrected ATR spectra closely resemble the acridine transmission spectrum over the entire wavelength range.This, in itself, is a critical test for the above approach and implies that the size of the scattering sites in the adsorbed layer is somewhere between 10 and 200 nm, a result which is in accord with independent determi- nations of cluster sizes in similar polymers to EPI-26.46 Results and Discussion Kinetic Curves Fig. 6(a) shows a plot of the ATR absorbance, AATR(A, O), at the isosbestic wavelength (A = 253 nm) and at a fixed angle of incidence (0 = 67.34"),measured as a function of time after a 250 ppm aqueous solution of polyelectrolyte EPI-26 (adjusted to pH 3.0) was introduced into the ATR cell holder containing a silica waveguide. Fig. qb) shows the correspond- ing plot of the scatter-corrected ATR absorbance, AA,R(A, 0),,,, , calculated via the procedures described above. The shape of these curves, which is typical for the adsorption of this polyelectrolyte onto silica under a wide range of solution conditions, clearly illustrates the extremely slow kinetics of J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 1.Ol 0.0b 5 1'0 1.5 20 25 time/h Fig. 6 ATR absorbance, A,&, 0), us. time for the adsorption of the EPI-26 polyelectrolyte at the silica/aqueous solution interface. (a) and (b), respectively, represent the raw and scatter-corrected ATR absorbance, measured at an angle of incidence 8 = 67.34" and at the isosbestic wavelength, 1 = 253 nm, after a 250 ppm EPI-26 poly-electrolyte solution (adjusted to a pH of 3.0 and a salt concentration of 0.01 mol dm-3 KCI) was introduced into the ATR cell at ambient temperature.the adsorption process, both of these curves only begin to show a plateau region after 15 h adsorption time. Given that the magnitude of AA,R(A, 0),at any fixed 8 and A, depends on the surface excess and conformation of the adsorbed EPI-26 polyelectrolyte, this would be expected to remain constant with time once equilibrium was reached. The gradual increase in AATR(A, 0) in Fig. 6 thus represents an approach to equi- librium either via a slow conformational change of the adsorbed polyelectrolyte, resulting in more chromophoric groups located nearer to the reflecting interface, via the adsorption of more polyelectrolyte or via some combination of both of these processes.Fig. 7 shows the effect of sudden changes of solution condi- tions (e.g.changes in pH or ionic strength) on the kinetics for the polyelectrolyte adsorption process. The diagram is a plot of AAT,(A, 0), collected for fixed O = 67.34' and I = 253 nm, 0s. time after a 100 ppm aqueous solution of polyelectrolyte EPI-26 (initially adjusted to pH 3.0) was introduced into the ATR cell holder [Fig. 7(a)]. Fig. 7(b) shows the effect of a change in ambient pH from 3.0 to 4.1, Fig. 7(c) that due to addition of potassium chloride salt, resulting in a concentra- tion of 0.01 mol dm-3 KCl, Fig. 7(d) a change of pH from 4.1 -2.5 a2.0-* x 0.0 v 0 32 64 96 128 160 time/h Fig. 7 ATR absorbance, A,,(253 nm, 67.34"),us.time collected for O = 67.34" and 1 = 253 nm after a 100 ppm aqueous solution of polyelectrolyte EPI-26, initially adjusted to pH 3.0 (a), was intro- duced into the ATR cell holder. (b)-(e) Sudden changes in pH or ionic strength of solution: (b)change to pH 4, (c) addition of 0.01 mol dm-3 KCl, (6)change to pH 5.2, (e)change to pH 2.3. cf)Rinse with triply distilled water and (9)rinse with HCI acid (pH 1.6). to 5.2 and Fig. 7(e)further change of pH from 5.2 to 2.3. Fig. 70 and (g), respectively, correspond to rinsing the ATR cell with the Milli-Q water and with dilute hydrochloric acid (pH 1.6). All of these changes in solution conditions clearly have a significant effect on the magnitude of AATR(A, 0) and thus a significant effect on the adsorbed polyelectrolyte conforma- tion.The increase in A,&, 0) with increasing pH and ionic strength [i.e. Fig. 7(a)-(e)]may be understood as an increase in the surface excess of the polymer which occurs as a result of two factors: (i) fewer PH+ *moieties on the polymer back- bone, which leads to a decrease in intra- and inter-molecular repulsion between polymer strands and hence allows a denser packing of polymer on the surface; and (ii) an increase in the (negative) surface charge of silica with pH (the isoelectric point of silica is slightly less than 347-49), which will lead to an increased electrostatic attraction of the positively charged polyelectrolyte with the surface. Conversely, the dramatic drop in AATR(I, 0) which occurs when the pH is lowered to 2.3 [Fig.7(e)] may be attributed to an expansion and partial desorption of the polyelectrolyte layer, both of which occur as a result of a positive surface charge at this pH and a large number of PH' moieities on the polymer backbone. The desorption process continues during the rinsing stages [Fig. 7cf) and (g)].The kinetics of both the adsorption and desorp- tion phases of Fig. 7 are extremely slow, at least 10 h are required for a plateau region to be reached after any solution change. Close inspection of these plateau regions also reveals slow changes in AATR(I, 0) even after extremely long (>92 h) equilibration times. This suggests that equilibrium conforma- tions for these systems may never be achieved.Adsorbed Conformation and Charge Distribution At any adsorption time, t, in Fig. 7, the conformation and charge distribution of the adsorbed polyelectrolyte layer may be determined by collecting variable-angle ATR spectra and analysing these via the procedure described above. An example of this, for t = 1 h, is shown in Fig. 8. Fig. 8(a) and (b),respectively, show the raw and scatter-corrected variable- angle ATR spectra and Fig. 8(c) shows the resulting Fi(f)us. [ plots, for the P and PH+ species, obtained by carrying out the required linear regression analyses. To minimize the effect of a small amount of solvatochromic band broadening which occurs in the ATR spectra for the acridine chromophore, a relatively narrow wavelength region (A = 247-260 nm) was chosen for this fit [i.e.for the linear regression involving eqn.(1111. Because of this narrow wavelength region, only the A [=Fp(f)] and B [=F,,+( f)] parameters in eqn. (11) could be determined with any degree of significance. The C and D parameters could not be determined. The moment analysis of these data was therefore carried out solely via linear regres- sion analysis involving eqn. (13). Fig. 8(c) shows the result of a linear regression fit of the F,([) and FpH+(f)data using two terms (moments) in eqn. (13). The curvature expressed in both sets of data clearly requires at least two moments to be used in this equation to obtain a good fit. This illustrates the fact that the variable-angle ATR experiment is sensing the spatial distribution of the P and PH' chromophores with respect to the interface, as well as their surface excesses. The present level of noise in the data, however, currently prevents the determination of more than two moments.Experimental modifications to the apparatus to reduce the level of noise, in anticipation of determining higher moments of the chromo- phore distributions, are currently under way. In principle, there is no reason why such higher moments could not be extracted from this type of experiment in the future. The results for the fit shown in Fig:8(c) are the following: rp = (0.33f0.01) x rnol dm-2, Zp = 13.7 & 1.5 nm, J. CHEM.SOC. FARADAY TRANS., 1994, VOL. 90 1257 excess of the polymer [l-= rp+ rPH+= (0.89 & 0.03) x mol dm-2 = 0.87 f0.03 mg m-'1 and the degree of ioniza-tion of the adsorbed polyelectrolyte [asurf= rPH+/(rP+ rpH+)= 0.631.Given that the acridine chromophore is distributed randomly along the EPI-26 polyelectrolyte back- bone, the mean thickness of the adsorbed polyelectrolyte, Z, may be calculated from: z = asurf ZPH + + (1 -asurf)zP (22) u.u .,,.,III 2!30 240 250 260 270 At t = 1 h, the polyelectrolyte thus has a mean thickness of wavelengthlnm 7.7 k 1.7 nm. This value may also be determined indepen- dently by performing a single-chromophore, single-wavelength fit of the AATR(A, 0) data at the isosbestic wavelength. It was found, however, that more accurate values 0.4 for Z and r could be obtained via the above multiple- A m wavelength fit.5 This type of analysis may now be carried out for any sur- c rounding solution conditions for the adsorbed polyelectrolyte q* 0.2 and at any time during the adsorption process. Fig. 9 shows other examples of Fi(5) data collected at various adsorption times, and for a variety of solution conditions, during the 0.0 230 240 250 260 270 experimen_t shoyn in Fig. 7. The curvature expressed in all of wavelengthlnm these Fi(Q us. 5 plots again clearly demonstrates that the variable-angle ATR experiment is sensitive to the adsorbed chromophore conformation and that at least two moments 0.~~1~ E 0.451 must be used in the moment equation [eqn. (13)] in order to obtain a good fit to the data. Moreover, changes in the degree of curvature between data collected under different T1:I0.35 solution conditions [i.e.differences between the plots in Fig. 9(a), (b) and (c)] suggest that the adsorbed layer takes on different conformations depending on the surround- ing solution conditions. Table 1 shows the results (calculated 0.15 r,Z,, ZPH+ and asurfvalues) for the variable-angle ATR data kI 40 60 80 100 120 shown in Fig. 9, as well as for other solution conditions and unm adsorption times for the experiment described in Fig. 7. Sig-Fig. 8 Raw (a) and scatter-corrected (b) variable-angle ATR (p nificant variations in the r,Z,, ZPH+ and asurfvalues listed in polarized) spectra collected at t = 1 h during the experiment Table 1 clearly illustrate that the variable-angle ATR experi- described in Fig.7.(c) Resulting Fit[) 0s. f plots, for I, P and 11, PH+ ment is sensitive to subtle changes in the structure and charge species, obtained by carrying out the linear regression analysis involving eqn. (1 1). The solid lines in (c) represent the best fits obtain- distribution of the adsorbed polyelectrolyte which are able using two terms (mo-ments) in eqn. (13)for the final linear regres- induced by changing the surrounding solution conditions. All (0)FP(5).Results of this fit are shown in of the variations in adsorbed conformation and charge dis- sion analysis: (0)FPH+(<), Table 1. tribution shown in Table 1 agree well with the qualitative description used in the previous section to describe the rpH+ = (0.56 f0.02) x mol dm-2 and :pH+ = 4.2 1.9 adsorption kinetics.For example, a change in ambient pH nm. These values give information about the conformation, from 5.2 (t = 66 h) to 2.3 (t = 120 h) leads to partial desorp- degree of charge and charge distributon of the adsorbed tion (rbecomes significantly smaller) and expansion (2, and layer: the Z values show that the positively charged PH' seg-ZPH+ become significantly larger) of the adsorbed poly- ments of the polymer are on average closer to the silica electrolyte layer. This result may be explained by the fact that surface than the neutral P segments (this is not terribly sur- at pH 2.3 the degree of ionization of the adsorbed layer is prising given that at pH 3.0 one would expect a fully larger than at pH 5.2 and also the silica surface will have a hydroxylated silica surface to be slightly negatively positive surface charge at pH 2.3 and a negative surface The values of Ti give both the total surface charge at pH 5.2.47,48 Both of these effects will lead to an Table 1 Calculated valuesa of r (polymer surface excess), Z, (mean separation distance of P chromatographores), Z,,, (mean separation distance of PH chromophores) and asurf(degree of ionization of the adsorbed polyelectrolyte) from variable-angle ATR absorbance spectra + collected at various adsorption times for the experiment described in Fig.7 time/% solution conditions ~~ r/mg m - Z,/nm Gi+/nm %urf 1 24 29.5 48 66 120 142 pH 3.0 pH 3.0 pH 4.1 pH 4.1, 0.01 mol dm-3 KCI pH 5.2, 0.01 mol dm-3 pH 2.3,0.01 mol dm-3 KCI rinse with acid (pH 1.6) rinse with acid for one week 0.87 (k0.03) 0.89 (f0.02) 1.09( f0.03)1.89 ( f0.06) 2.26 (kO.11) 0.61 (k0.02) 0.46 (f0.02) 0.29 (fO.01) 13.7(k1.5) 11.8 (f0.9) 8.4(f1.4)14.1(21.8) 5.0( f2.4) 16.5(k1.1)19.0(k3.5) 21.8(k3.1) 4.2 ( f1.9) 8.5 (& 1.0) 4.4 (* 1.2)8.1 (f1.5) 5.4( f2.5)10.7 (f1.7) 15.4 (f 1.7) 17.9 (f2.2) 0.63(k0.04)0.67(kO.04) 0.65 (kO.04) 0.63(k0.05)0.62(k0.05) 0.74( f0.05) 0.73(k0.05)0.71(kO.05) a These data were originally presented in an earlier p~blication~~ without the evanescent wave scattering correction.All numbers in Table 1 are corrected for scattering and represent absolute values. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 u 0.4875ill 0 7g,2-0*3750/ NI & 0.60 ? 0.45“I 0 .1 5 1 40 60 80 100 120 4/n i-o.800t0 40 60 80 100 120 t/nm Fig. 9 Fit) us. t plots calculated from variable-angle ATR spectra collected at adsorption times t = 24 (a),29.5 (b)and 48 h (c) during the experiment described in Fig. 7. The solid lines represent the best fits obtainable using two moments in eqn. (13), symbols as in Fig. 8. The resulting r,Z,, Z,,, and values calculated from these fits are shown in Table 1. expansion and partial desorption of the adsorbed layer, first by increasing the electrostatic repulsion between the polymer strands and secondly by introducing an electrostatic repul- sion between the adsorbed (positively charged) poly- electrolyte and the positive surface. The results in Table 1 also show that rinsing the adsorbed polyelectrolyte layer with dilute hydrochloric acid (pH 1.6) leads to further expansion and desorption, a process which is not complete even after a week of rinsing. This result suggests that adsorption of the EPI-26 polyelectrolyte at the silica/aqueous solution interface is effectively an irreversible process. Degree of Ionization Another interesting feature of the adsorbed polyelectrolyte layer is the fact that the degree of ionization of the poly- electrolyte at the surface, Q,,~, is nearly always different to ~~that of the polyelectrolyte in solution, Q~ An example of this is shown in Fig.10, where a direct comparison can be made between the variable-angle ATR spectra of the surface adsorbed polyelectrolyte [Fig.lqa)] and the transmission spectrum of the surrounding polyelectrolyte solution [Fig. lqb)], collected simultaneously at an adsorption time of t = 1 h for the experiment described in Fig. 7. The difference in the relative peak heights of the P and PH’ species for the two diagrams immediately shows that the degree of ioniza- tion for the polyelectrolyte at the surface is less than that for the polyelectrolyte in solution. Quantitative analysis of these I wavelengt h/nm 1.6 0.0 230 240 250 260 270 wavelengthlnm Fig. 10 Comparison between the variable-angle ATR spectra of the surface-adsorbed EPI-26 polyelectrolyte (a) and the transmission spectrum of the EPI-26 polyelectrolyte solution (b),collected at an adsorption time oft = 1 h during the experiment described in Fig.7. The shape of these spectra clearly shows that the degree of ionization of the polyelectrolyte adsorbed to the surface, aSud= 0.63, is signifi- cantly different from that in solution, amlution0.78, I, P; 11, PH+.= spectra reveals = 0.63 (from Table 1) and a,lution= 0.78. Table 2 lists values of a,,,f and asolutiondetermined, in an iden- tical manner, for a variety of solution conditions during the experiment described in Fig. 7. The large differences between a,,,f and asolution,which are seen for all of the solution condi- tions listed in Table 2, illustrate the significant effect that the presence of the silica surface has on the pK, of the ionizable P groups attached to the adsorbed polyelectrolyte backbone.The pK, of such groups will be affected by the local field resulting from charges present at the silica/water interface, as well as charges on the polyelectrolyte backbone, and would be expected to be significantly different to that of P groups in bulk solution. The insensitivity of a,,,, to pH and ionic strength, however, suggests that the hydrophobic aggregates present in the adsorbed layer may be inhibiting the ionization mechanism of the polyelectrolyte layer. These aggregates could ‘bury’ ionizable moieities inside hydrophobic environ- ments and thus distort the ionization equilibrium of the surface-adsorbed polyelectrolyte. Such effects have been observed previously in other interfacial systems, e.g.ref. 49. Conclusion We have demonstrated the potential of a simple evanescent wave spectroscopic technique for measuring concentration profile information, p(z), for multiple chromophores attached to the backbone of an adsorbed polymer. To begin with, we ~~~~~. Table 2 Experimentally determined values of aSur,and .a,,u!ion for various adsorption times during the experiment described in Fig. 7 time/h solution conditions ‘surf ‘solution 1 pH 3.0 0.63 (kO.04) 0.78 (kO.01) 24 pH 3.0 0.67 (kO.04) 0.78 (kO.01) 29.5 pH 4.1 0.65 ( f0.04) 0.47 ( f 0.01) 48 pH 4.1, 0.01 mol dm-3 KCI 0.63 (k0.05) 0.48 (kO.01) 66 pH 5.2, 0.01 mol dm-3 KCl 0.62 ( k0.05) 0.15 (f 0.01) 120 pH 2.3, 0.01 mol dm -3 KCI 0.74 ( k0.05) 0.98 ( f0.01) J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 have shown that the variable angle of incidence ATR tech- nique can be used to determine rigorously (without any assumptions about the structure of the polymer layer) the surface excess, r,the degree of ionization, asurf,and the mean separation distance from the interface, Z, of charged and uncharged segments attacked to the backbone of a model polyelectrolyte. The kinetics of polyelectrolyte adsorption (and desorption) can be easily followed uia this technique and the experiment can be performed for any surrounding solu- tion conditions. Given the present level of noise in the experi- mental apparatus, only the first two moments of the p(z) function may be determined.With improvement of the apparatus, to reduce the level of experimental noise and to improve sensitivity towards chromophores with small molar absorption coefficients, it is anticipated that this technique will be able to provide even more information about model adsorbed polyelectrolytes (such as EPI-26), as well as improve the ability to study adsorbed polymers which have not been specifically ‘tagged’ with a high molar absorption coefficient chromophore. We would like to thank Mr. Rodney Parr (ICI Research Group) and Dr. San Thang (CSIRO Division of Chemical and Polymers) for their tremendous help with the synthesis of the model polyelectrolyte. Mr. Richard Mathys is thanked for his excellent design ideas and for construction of a large portion of the apparatus.M.T. acknowledges the receipt of an Australian Postgraduate Research Award and an ICI (Australia) Postgraduate Scholarship. Financial assistance from the Australian Research Council, the Advanced Mineral Products Centre (University of Melbourne) and ICI (Australia) Specialty Chemicals Group is also gratefully acknowledged. References 1 R. W. Armstrong and U. P. Strauss, in Encyclopedia of Polymer Science and Technology, ed. H. F. Mark and N. G. Gaylord, Wiley, New York, 1969, vol. 10, pp. 781-861. 2 J. Lyklema, in Modern Trends of Colloid Science in Chemistry and Biology, ed. H. F. Eicke, Birkhauser Verlag, Basel, 1985, pp. 55-73. 3 F. Macritchie, Ado. Protein Chem., 1978,32, 283.4 J. D. Andrade and V. Hlady, Adv. Polym. Sci., 1986,79, 1. 5 (a) A. Takahashi and M. A. Kawaguchi, Adv. Polym. Sci., 1982, 46, 1;(b)Ado. Colloid Interface Sci., 1992,37, 219. 6 M. A. Cohen Stuart, T. Cosgrove and B. Vincent, Ado. Colloid Interface Sci., 1986,24, 143. 7 S. Sugai and G. Ebert, Adv. Colloid Interface Sci., 1986,24, 247. 8 G. J. Fleer, and J. Lyklema, in Adsorption from Solution at the SolidlLiquid Interface, ed. G. D. Parfitt and C. H. Rochester, Academic Press, London, pp. 153-218. 9 T. Cosgrove, J. Chem. SOC., Faraday Trans., 1990,86, 1323. 10 T. Cosgrove and B. Vincent, in Fluid Interfacial Phenomena, ed. C. A. Croxton, Wiley, New York, 1986. 11 H. A. Van der Schee and J. Lyklema, J. Phys. Chem., 1984, 88, 666 1.12 J. Papenhuijzen, H. A. Van der Schee and G. J. Fleer, J. Colloid Interface Sci., 1985, 104, 540. 13 0.A. Evers, G. J. Fleer, J. M. H. M. Scheutjens and J. Lyklema, J. Colloid Interface Sci., 1985, 111, 446. 14 M. R. Bohmer, 0. A. Evers and J. M. H. M. Scheutjens, Macro-molecules, 1990, 23, 2288. 15 M. K. Granfeldt, S. J. Miklavic, S. Marcelja and C. E. Wood-ward, Macromolecules, 1990,23,4760. 16 M. K. Granfeldt, B. Jonsson and C. E. Woodward, J. Phys. Chem., 1991,95,4819. 17 T. Cosgrove, T. M. Obey and B. Vincent, J. Colloid Interface Sci., 1986, 111, 409. 18 T. Cosgrove, J. S. Phipps and R. H. Richardson, in Surface X-Ray and Neutron Scattering, Springer Proceedings in Physics, ed. H. I. Zabel and I. K. Robinson, Springer Verlag, Berlin, 1992, vol.61. 19 M. Kawaguchi, K. Hayashi and A. Takahashi, Colloids Su$, 1988,31, 73. 20 M. Malmsten and B. Lindman, Langmuir, 1990,6,357. 21 T. Cosgrove and K. Ryan, Langmuir, 1990,6,136. 22 K. K. Fox, I. D. Robb and R. J. Smith, J. Chem. SOC., Faraday Trans. I, 1974, 70, 1186. 23 M. C. Cafe and I. D. Robb, J. Colloid Interface Sci., 1982, 86, 411. 24 P. A. Williams, R. Harrop and I. D. Robb, J. Chem. SOC., Faraday Trans. 1, 1985,81,2635. 25 I. D. Robb and M. Sharples, J. Colloid Interface Sci., 1982, 89, 301. 26 P. L. Gramain and Ph. Myard, Macromolecules, 1981,14, 180. 27 T. Cosgrove, T. L. Crowley, B. Vincent, K. G. Barnett and Th. F. Tadros, Faraday Symp. Chem. SOC., 1981,16,101. 28 D. E. Graham and M. C. Phillips, J.Colloid Interface Sci., 1979, 70,415. 29 J. E. Gebhardt and D. W. Fuerstenau, Colloids Surf., 1983, 7, 221. 30 R. P. Vedula and H. G. Spencer, Colloids Surf,1991,58,99. 31 J. Papenhuijzen, G. J. Fleer and B. H. Bijsterbosch, J. Colloid Interface Sci., 1985, 104, 553. 32 J. C. Charmet and P. G. de Gennes, J. Opt. SOC. Am., 1983, 73, 1777. 33 I. Caucheteux, M. Hervet, R. Jerome and F. Rondelez, J. Chem. SOC., Faraday Trans., 1990,86,1369. 34 (a)W. Lukosz and R. E. Kunz, J. Opt. SOC.Am., 1977, 67, 1607; (b) 1615. 35 W. Lukosz and R. E. Kunz, Opt. Commun., 1977,20,195. 36 H. Kuhn, J. Chem. Phys., 1970,53, 101. 37 P. Suci and V.Hlady, Colloids Surf., 1990,51,89. 38 G. Rumbles, A. J. Brown and D. Phillips, J. Chem. SOC.,Faraday Trans., 1991,87, 825. 39 M. Trau, F. Grieser, T. W. Healy and L. R. White, Langmuir, 1992,8, 2349. 40 C. L. McCormick and G. S. Chen, J. Polym. Sci., 1984,22,3633. 41 M. Trau, Ph.D. Thesis, University of Melbourne, 1992. 42 W. E. Wentworth, J. Chem. Educ., 1965,42,96. 43 G. A. Schumacher and T. van de Ven, Langmuir, 1991,7,2028. 44 M. Kerker, The Scattering of Light and Other Electromagnetic Radiation, Academic Press, New York, 1969. 45 P. N. Pusey, in Colloidal Dispersions, ed. J. W. Goodwin, Royal Society of Chemistry, London, 1982. 46 C. E. Flynn and J. W. Goodwin, in Polymers as Rheology Modi- fiers, ACS Ser. 462, ed. D. N. Schulz and J. E. Glass, American Chemical Society, Washington DC, 1991. 47 P. J. Scales, F. Grieser, T. W. Healy, L. R. White and D. Y. C. Chan, Langmuir, 1992,8,965. 48 R. 0.James, in Advances in Ceramics: Ceramic Powder Science, ed. G. L. Messing, K. S. Mazdiyashi, J. W. McCaulley and R.A. Haber, American Ceramic SOC., Westerville, OH, 1987, vol. 21. 49 B. Murray, J. Godfrey, F. Grieser T. W. Healy, B. Lovelock and P. J. Scales, Langmuir, 1991,7, 3057. Paper 3/05307E; Received 3rd September, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949001251

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(9), 1261-1263 Comparison of the Electrokinetic Properties of the Silica Surface Dave E. Dunstan Department of Chemical Engineering The University of Melbourne Parkville, Victoria 3052, Austra lia The electrokinetic potentials and surface charge densities of the silica surface are interpreted from both electro- phoretic mobility and electro-osmosis measurements. Comparison of the data shows smaller magnitudes for both the potentials and surface charge interpreted from the electrophoretic mobility over the range of KCI concentration studied (10-5-0.1 mot dm-3). The interpreted surface charge densities increase indicating that effective adsorption of the negative chloride ions to the negative surface occurs. The silica surface is therefore postulated to be non-classical in nature.It is suggested that the retardation effect and therefore the electro-phoretic potentials are underestimated in the electrophoresis problem. The electrokinetic equations describing colloidal phenomena are undergoing close experimental examination at present. There are several reasons for this. The continuum model pro- posed in the theory is of fundamental physical interest and the practical implications of the measurements and their interpretation are wide.' It is generally accepted that a test of the theory is that it yields the same electrokinetic potential from different measurements on the same surface. To date this has not been satisfactorily achieved. Theoretical modifi- cations have been made in an attempt to obtain consistency between the interpreted potential^.^,^ However, no com-prehensive study of a single system which enables vindication of the modified or fundamental theory has yet been made. This work presents and compares data obtained from two direct current (dc) electrokinetic measurements of the silica surface over a range of electrolyte concentrations.The elec- trophoretic mobilities of silica particles and the electro-osmotic mobilities of silica plates were measured over a range of KCl concentrations and the interpreted potentials com-pared.4 Further insight into the nature of the silica surface and the electrokinetic equations is obtained. Experimental KCl obtained from Fluka (99.5%) was recrystallized and baked at 600°C for 4 h.Deionized water was further purified by double distillation from a KOH-KMnO, first stage and stored in a Pyrex vessel until used. The water used was of surface tension 72.6 mN m- ',conductivity 0.9 mS cm- ' and pH 5.8 (1.5 x mol dmP3 H,CO,). All measurements were conducted at pH 5.8. Two sizes of silica particles were prepared using the method of Stober et a1.' to yield particle sizes of 50.4 and 165.0 nm radius. These particles are reported to be of lower density (2.0 g cm-3) than amorphous silica (2.2 g cm-3).6*7 Particle sizing was done using electron micros- copy on more than 300 particles. The silica particles were extensively cleaned by centrifugation-decantation ca. 100 times over a period of one year in Teflon centrifuge tubes.This ensured that all ammonia and other possible impurities were removed. Polystyrene lattices, used as tracers, were obtained from Interfacial Dynamics Corporation (sulfate variety, 300 nm in diameter). These were extensively cleaned by centrifugation-decantation 50 times to form an opalescent suspension in H20 at a volume fraction of 0.05 which showed no surface-active contamination to be present. All the reported electrokinetic measurements were made using a Coulter Delsa 440 Electrophoresis apparatus. The mobilities of the silica particles were measured as a series by addition of high concentrations of KC1 to a dilute stock solu- tion. Using this method the volume fraction of the suspension was kept approximately constant. The final concentration was determined by mass.The silica particle mobilities were measured at both stationary layers to ensure that no con- tamination artefacts were present during the measurement. Measurement of the electro-osmotic mobility required the use of a tracer particle, polystyrene latex, and measurement of the apparent mobility profiles. Latex particles were used as tracers in preference to silica particles as the latices give nar- rower, less noisy mobility peaks. To evaluate the electro- osmotic mobility of the silica cell wall from the apparent mobility profile in the Delsa cell (aspect ratio 3.25) the Komogata equation was used.' The Komogata equation describes the relationship between the apparent particle mobility, U,, , the electrophoretic mobility, U, , and the electro-osmotic mobility, U,, ,as a function of position in the cell.U,, = U, -O.86Ue0+ 1.86Ue0~/h)2 where y/h is the fractional distance from the centre of the cell to the wall. The symmetry of the parabolic profiles indicated that the cell was free of contamination. The electro-osmotic potentials were interpreted from the electro-osmotic mobilities of the latex tracer particles using the Smoluchowski equation.' This analysis was consistent with that of O'Brien and White4 for infinite particle radius. Results and Discussion The electrokinetic potentials and surface charge densities interpreted from the measurements are shown in Fig. 1 and 2. Both the electro-osmosis and electrophoresis measurements are interpreted using the theory of O'Brien and White.4 In the limit of high KU the theory turns to the well-known Smol- uchowski result.Here, K is the reciprocal Debye length and a the particle radius. Several trends are immediately obvious. The potentials and charge densities interpreted from the electro-osmosis mea-surements are greater in magnitude than those interpreted from the electrophoretic mobilities of the particles over the range of electrolyte concentration employed. The surface charge densities increase with increasing electrolyte concen- tration for the electro-osmotic measurements. The surface charge densities interpreted from the electrophoresis mea- surements on the particles show a slight maximum at ca.0.01 mol dm-3 KCl. Both data sets indicate that negative ions are being adsorbed to the negative silica surface. [KCl]/mol dm-3 Fig. 1 Interpreted electrokinetic potentials of the silica surface us. KCl concentration. The curves are best fits to the data and are shown to help distinguish the different data sets. The derived electro- phoretic potentials for the two different particle sizes are: (+) 504 nm particles and (0)1650 nm particles. (W) Electro-osmotic and (0) streaming potential" data. The possibility that experimental artefacts are present must be discussed before a physical interpretation of the data is made. The particles are deemed to be This would imply that the surface area of the particles is larger than expected.Therefore, the driving force on the particles should be larger and the interpreted potentials and surface charge densities should therefore be higher than those interpreted for the flat plates. The derived potentials and surface charge den- sities for the particles of two different sizes would also not be the same, as their surface areas are different. Essentially the surface-to-bulk ratio will be different for the two particles and thus any porosity should not give the same contribution to the mobility for the different particles. The possibility that the particle surfaces are porous or that contamination is present does not account for the increasing electrokinetic N E Q)P r 0 -2 Q) f 10-6 lo4 lo-' loo [KCl]/mol dm-3 Fig.2 Interpreted electrokinetic surface charge densities of the silica surface us. KCl concentration. (+) 504 and (D) 1650 nm particles. (0)Values interpreted from the electro-osmosis measurements. The curves are drawn to guide the eye. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 surface charge density with electrolyte concentration. Both the particles and plates are made of silica and were cleaned stringently, as discussed in the Experimental. As such, both should behave as silica exposed to electrolyte. To ensure that the electro-osmotic measurements were interpreted correctly using the Komogata equation to describe the hydrodynamics of the Delsa cell, data obtained by other workers for silica plates in streaming potential measurements are also shown in Fig.1 (0)''The data for the two sets of independent mea- surements agree within experimental error, indicating that the cell walls are free of contamination and that the Komo- gata equation is correct for the aspect ratio of the cell. The streaming potential measurements are for an effectively infin- ite aspect ratio. Given that the experimental data represent a true compari- son of the silica surface measured by several techniques, let us return to a physical discussion of the data. If the theory was a correct physical description of the measured phenomena, the interpreted potentials would be the same. The physical reason for the occurrence of both phenomena, electrophoresis and electro-osmosis, is the finite spatial separation of charge which occurs at the electrolyte/solid interface.The plane of shear in both systems defines the elec- trokinetic potential. The spatial distribution of ions and the fluid flow in the electrolyte, outside the shear plane, are described by the continuum equations, the Poisson-Boltzmann and Navier-Stokes equations, re~pectively.~ The usual assumptions of a continuous dielectric, constant vis- cosity solvent, with point charges for the ions, apply. In the electro-osmotic measurements, electro-osmotic mobility is induced in the solvent under the influence of the applied field by the flow of the asymmetric charge that resides in the inter- facial region. The flow of charges carries liquid with it.This is a contradiction of the continuum model. The point charges should have zero interaction with the solvent and therefore induce no flow. Point charges would also have infinite con- ductivity. This is accounted for in the theory by using the measured values of the specific ionic conductivities. The electrophoresis problem is of a similar nature at high electrolyte concentration. In the Smoluchowski limit, the relative motion of the particle and electrolyte is induced by the electro-osmotic motion around the particle. At lower elec- trolyte concentrations the particle moves due to the force exerted on it by the electric field and the electrokinetic surface charge. The counterion charge in the diffuse double layer is distorted i.n the electric field and gives rise to the retardation effect.The dominant retardation effect on the particle mobility is due to the fact that the electric field seen by the particle is less than the applied field combined with an extra viscous drag due to the diffuse double layer being carried with the particle. This polarization of the diffuse double layer around the particle is responsible for the very large dielectric response observed for colloidal particles.' From the above discussion it is apparent that one reason for the differences in the interpreted potentials and surface charge densities is that the retardation effect is incorrectly calculated by the theory. The spatial distribution of ions in the double layer is necessary to calculate the retardation effect in the electrophoresis problem, while in the electro- osmosis problem the exact spatial distribution of the charge need not be known. It is therefore readily concluded that the Poisson-Boltzmann equation is not an accurate description of the diffuse layer in this case.This interpretation is consis- tent with dielectric response measurements on colloidal silica which show the theory to underestimate the dielectric response significantly. l2 Essentially, the experimentally mea- sured dipole moments are much larger than could be expected reasonably from the theory. Therefore, the interpre- J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 tation of the electrophoresis measurement underestimates the electrokinetic potential. Consistent with the above discussion is the observed increase in magnitude of the electrokinetic surface charge with electrolyte concentration.The electrokinetic surface charge should decrease with electrolyte concentration as positive ions are adsorbed to the negative surface as the bulk electrolyte concentration increases. The data indicates that negative ions are adsorbed to the negative surface as the con- centration increases. This is non-classical behaviour and further evidence that the Poisson-Boltzmann description of the interfacial charge distribution is not complete. A classical mechanism for the adsorption of negative ions to a negative surface has yet to be proposed. Recent measurements of the electrokinetic properties of hydrocarbon particles in an elec- trolyte have shown the presence of finite mobilities, which are interpreted as resulting from changes in the chemical poten- tial of the solvent due to the presence of an interface.l39l4 While the silica surface is intrinsically charged oia the disso- ciation of hydroxy groups at the measurement pH, the surface may still cause a perturbation in the interfacial water.As such, the interfacial region may be described by a com- bination of both Poisson-Boltzmann electrostatics and a spa- tially varying solvent chemical potential. Note that recent measurements on the interaction forces between silica surfaces have shown the continuum description to be remarkably good down to small separation distances of the order of several nanometres.” This may be reconciled with the data presented here by assuming that the classical electrostatic interaction is observed at small separation dis- tances.It is hoped that future work on different electrolytes will further elucidate these interesting phenomena. References 1 R. J. Hunter, The Zeta Potential of Colloid Science, Acadmic Press, New York, 1987. 2 E. H. B. de Lacey, Ph.D. Thesis, Australian National University, 1982. 3 C. S. Mangelsdorf and L. R. White, J. Chem. SOC., Faraday Trans., 1990,86,2859. 4 R. W. O’Brien and L.R. White, J. Chem. SOC.,Faraday Trans. 2, 1978,74, 1607. 5 W. Stober, A. Fink and E. Bohn, J. Colloid Interface Sci., 1968, 26, 62. 6 R. K. Iler, The Chemistry of Silica, Wiley, New York, 1979. 7 G. H. Bogush, M. A. Tracy and C. F. Zukoski, J. Non-Crystalline Soids, 1988, 104, 95. 8 S. Komogata, Res. Electrotech. Lab. Tokyo, Comm. No. 1933, 348. 9 M. Von Schmoluchowski, Z. Phys. Chem., 1918,92, 129. 10 P. J. Scales, F. Grieser and T. W. Healy, Langmuir, 1990,6, 582. 11 E. H. B. de Lacey and L. R. White, J. Chem. SOC., Faraday Trans., 1981, 77, 2001. 12 L. A. Rosen and D. A. Saville, Langmuir, 1991,7,32. 13 D. E. Dunstan and D. A. Saville, J. Chem. SOC.,Faraday Trans., 1992,88,2031. 14 D. E. Dunstan and D. A. Saville, J. Chem. SOC., Faraday Trans., 1993,89, 527. 15 A. Grabbe and R. G. Horn, J. Colloid Inter$ace Sci., 1993, 157, 375. Paper 3/07106E; Received 1st December, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949001261

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(9), 1265-1269 Vibrational Spectroscopic Analysis of Group 6 Metal Hexacarbonyls in the Solid State Upali A. Jayasooriya School of Chemical Sciences, University of East Anglia, Norwich, UK NR4 7TJ Vibrational spectra of group 6 metal hexacarbonyls in the solid state are analysed using the ‘oriented gas’ and ‘ latent symmetry’ approaches. An explanation of lifting of degeneracies and relative intensities is provided, including the intensity predictions for the factor group components of distortion-induced intensities of molecu- larly forbidden modes. Vibrational spectroscopy is a tecnique which is used rou-tinely to obtain structural information on molecules in phases where the intermolecular interactions present are neg- ligible.However, the same cannot be said about the solid state when there is appreciable intermolecular coupling. One reason for this is a dearth of simple crystal structure-spectra correlations, with regard to solid-state effects. When a highly symmetric molecule crystallises it com-monly adopts a site which has significantly lower symmetry than that of the molecule itself. In such cases it is always a matter of interest to investigate the extent to which the packing arrangement perturbs the molecular properties. This problem is of particular interest for species of 0,symmetry because there are Wigner-Seitz unit cells of 0,symmetry, the octahedron itself is not one of these. Octahedral molecules, then, are likely to be forced to adopt sites of rather low sym- metry.An excellent example is provided by the octahedral metal carbonyls M(CO),, M = Cr, Mo, W, all of which crys- tallise in Pnmu (D::) with four molecules per unit cell, each molecule occupying a site of C,symmetry.’ The first vibrational spectroscopic data on these were published in 1955 by Hawkins et a2.’ and this was followed by a series of studies mainly aimed at recognising and assign- ing the molecular vibrational modes.3 In 1973 single-crystal Raman experiments were used4 to obtain symmetry species of the vibrational modes in the solid state. Then in 1978, Kariuki and Kettle’ used Raman spectra of mixed crystals to show the presence of strong intermolecular vibrational coup- ling in the v3(CO), e, mode and the absence of such coupling effects in the case of the v,(CO),a, mode.Further, the weak bands at 1964 (Cr), 1966 (Mo) and 1957 (W) cm-l were iden- tified as Raman activity shown by the only infrared-active v6(CO), t,, mode of the isolated molecule. In the same year Scheuermann and Nakamoto, reported Raman spectra of chromium hexacarbonyl isolated in an argon matrix at 13 K and in the solid state at 292 and 13 K. These workers have observed many more bands in the solid-state spectrum at low temperature and have assigned frequencies to six inactive fundamentals of t,, ,t,, and t,, symmetries, which were pre- viously only inferred from overtone and combination modes. However, surprisingly, their assignments are based on an earlier crystal structure determination7 which has since been shown to be in error.’ All extant single-crystal vibrational spectroscopic data, consisting of Raman and reflection IR spectra, are due to Adams and co-worker~.~.~ The following is an attempt to understand some of the solid-state effects shown by these compounds. The ‘oriented gas model’g is first used to explain the intensity distribution in the carbonyl stretching frequency region of the Raman spectrum.Its applicability is illustrated using the v(CO), e, mode, which has been shown previously to be the more vibrationally coupled of the two Raman-ative v(C0) molecu- lar modes. Secondly the ‘latent symmetry approach’” is used to obtain qualitative intensity predictions on the whole of the IR and Raman spectra. In this technique, purely symmetry- based arguments are used to investigate the evolution of the real crystal structure from the isolated molecule via the latent or near-symmetries present in the solid state.During this process of evolution, any vibration allowed in the final factor- group analysis but forbidden under the molecular and/or latent symmetry (i.e. the corresponding matrix element =0, under these symmetries) is predicted to have a value for its matrix element which is proportional to the extent of devi- ation of the real structure from the latter symmetries. There- fore the useful latent symmetries are those involving the minimum of distortion from the real crystal structure. The utility of the latter approach is illustrated by providing expla- nations of the following experimental observations.The close similarity between the single-crystal Raman spectra of b,, and b,, factor-group symmetries.* Also similar to each other are the IR spectra of b,, and b,, factor-group symmetries. These observations are highlighted in the splitting pattern of the four factor-group components arising from the molecular e, carbonyl stretching mode. Here the b,, and b2g com- ponents are almost degenerate while a, and b3, ar.: substan-tially split.’ Further, molecular IR and Raman activities are almost strictly maintained through to the solid state so that only about half the factor-group predictions are realised in practice. The only exception to this is the molecularly Raman-inactive t ,, mode found in the carbonyl stretching region of the Raman spectrum.This shows up in the single- crystal Raman spectra with intensity only in the a, and b3, symmetries even though six Raman-active factor-group modes are predicted theoretically (Table 1). Discussion X-Ray Crystal Structure The complete crystal and molecular structure has been reported for chromium hexacarbonyl and therefore this compound is used mainly for the following discussion. However, the conclusions are of general validity for this whole class of isomorphous compounds. A careful examination of the crystal structure shows that these molecules crystallise with molecular three-fold axes approximately parallel to each other and to the crystallo- graphic ‘a’ axis [Fig. l(b)].This near-symmetry would not show up in the space group and thus will have no effect on the factor-group predictions, conventionally used to explain vibrational spectra in the solid state. Note that in this case, even if the molecular three-fold axes are exactly parallel, it would still not change the space group and hence the factor group. However, it is possible to use this near-symmetry or J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 1 Correlation between the factor group, site group, the most significant latent factor group, latent site group and the molecular point group symmetry of Cr(CO), latent latent factor molecular site factor group group molecular Oh D,, '2h D4h 0:: Oh / The effect of the tetragonal latent space group which is of only intermolecular origin, is indicated by dotted lines. This shows that when the degenerate sets of functions (xz, yz); elg and (x2 -y2, xy):,eZg under D6h are correlated with those in D,, , only the former degeneracy is maintained.'latent symmetry'" in order to simplify the application of the 'oriented gas approximation '' for the estimation of intensities of vibrational transitions in these solids. The Raman itensities of the factor goup components derived from the e, carbonyl stretching mode are chosen as an example. Oriented Gas Model The two tensors for the components of e, modes in the point group 0,are given by Poulet and Mathieu,' and are: -u 0 (1)[ : -; 2\] rrand -+a -These are to be expressed in terms of crystallographic axes using a similarity transformation relating the two axial systems, molecular and crystal.The relationship between these two axial systems is shown in Fig. 2, which gives the transformation matrix : The two tensor components of the molecular e, mode in terms of the crystal axes are therefore; Now the use of the projection operator technique in the factor group Diz gives the derived polarisability tensors cor- responding to the respective factor-group components as : 0 -J2a 0 O J2u:] 000 ab2g = ; ab3p = [o 0 a] (4)[ ;J2u 0 OUO These tensors when squared are proportional to the Raman intensities of the corresponding factor-group modes due to a single unit cell, or an aligned single crystal.If one were to compare the experimental single-crystal Raman spectroscopic data with these predictions, it is essential to take great care to compare spectra run under exactly similar conditions. Such data are not available at present. However, one may use the spectra from polycrystalline samples where all symmetries appear in the same spectrum. To do this, it is necessary to obtain expressions assuming all orientations of these unit cells with respect to the observer's axes with equal probabil- ity. Following the method given for this by Wilson et ~1.'~ one arrives at the following intensity ratio for these factor- group modes observable from a randomly oriented poly- crystalline sample. la, : Ibl, Ibzs Ib3, 1 2 2 : 1 (The v4 dependence of the intensities is assumed to be the same for all four modes because of the closeness of their 0 J2u 0 -J2u 0 absolute frequencies.) The polycrystalline Raman spectra of these hexacarbonyls [j2-; ;] and [-d: ;] (3) show that it is not possible to resolve all these factor-group J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 e Fig. 1 Crystal structure’ of Cr(CO), projected down the X-axis. (a) The two hexagonal sub-lattices, obtained by the molecular pairing (k, r) and (m,n), which combine to give the real structure. The molecules within a unit cell are named k, I, m,n as shown in Fig. 3. (b) The full structure showing the molecular three-fold axes parallel to each other and to the crystallographic X-axis.(c) The two approximately tetragonal sub-lattices, obtained by the molecular pairing (k, n) and (I, m),which combine to give the real structure. This tetragonal sym- metry is purely intermolecular in origin with no intramolecular counterpart. components in order to measure their intensities separately.’ However the b,, component which is assigned to ca. 2024 cm-’ (vide infra) stands well clear of the other three. There- fore the intensity ratios between this and the sum of the intensities of the a,, bl, and b,, modes are measurable and are ca. 1 : 6.2 and CQ. 1 :5.2 for Cr(CO), and Mo(CO),, Y I, 1 -1 -1X J3 J3 J3 1 -1 Y -1 -1 -Z J2 J* J2 0 I[: Y J6 J6 J3 Fig. 2 Relationship between the molecular (X, Y, 2)and crystallo- graphic (x, y, z) axes respectively.These results are in reasonable agreement with the predictions from the above calculation especially when one takes into account the approximate nature of the model used here. Much effort had been spent in the past in attempt- ing to improve the oriented gas model in general, particularly with reference to lattice modes.’ Latent Symmetry Approach There are many subtleties of the vibrational spectra of these systems not explained by the latter technique. Here we resort to the so-called ‘latent symmetry approach’,’’ where near- symmetries present in the solid state or those symmetries which are hidden from the space and factor-group decriptions, like the parallel three-fold axes in the present case, are made use of in the spectral interpretation.This is a qualitative approach providing valuable insight into the sym- metries of the significant perturbing potentials of the molecu- lar environment in the crystalline state. In order to make use formally of the fact that the molecu- lar three-fold symmetry axes are parallel in the solid state, one has to relate this latent symmetry element to the space- group symmetry of this material. The method of searching for the latent symmetries adopted here is to take one of the mol- ecules from the four symmetry-related molecules in the unit cell, and divide the total potential of its environment into the individual contributions made by each of the other molecules in the unit cell, together with all molecules translationally related to the latter.Fig. l(b) shows the crystal structure projected down the x axis and Fig. 1(a) shows the two hexagonal substructures combined to give the real structure [i.e. Fig. l(b)]. The mol- ecules within a unit cell are named k, l, m and n (Fig. 3) and the intermolecular potentials shown in Fig. l(a) are for the molecular pairs (k,I) and (m,n). These two substructures have clear hexagonal symmetry, and the three-fold axes in these hexagonal substructures are displaced with respect to each other along the z direction. If one were to bring these two structures to register by translations along z and x directions (necessary translations indicated in Fig. 3), the resulting structure will have the space group P6,lmcm; D&, with only two molecules per primitive unit cell in sites of symmetry D3d.The ‘distortions’ relating this structure and the real structure are purely non-primitive translations, and therefore to a first good approximation are not expected to affect the intensities of the internal mode vibrations. lo Fig. l(c) shows another possible pairing of these molecules, into (k,n) and (I, m)interactions. These are two substructures of approximately tetragonal symmetry. Here note that this tetragonal distribution is purely intermolecular in origin with no intramolecular counterpart. If one ignores the molecular symmetry in this case and consider only the intermolecular distribution, it approximates to the space group P4/mmm; Dih, with only one molecule per primitive unit cell in a site of D& symmetry.The only other possible pairing of molecules of (k,m) and (I, n) does not show any extra symmetry with each pair con- fined to a mirror plane already recognised in the real space group. In summary, only two latent-symmetry structures are therefore recognised with space groups P6,lmcrn; D&, and P4/mmm; Dih, with the former being the most significant with full intra- and inter-molecular symmetries. The symmetry correlations between the molecular point group 0,and the site and factor groups of latent and real space groups are given in Table 1. The four factor-group symmetries, active in the Raman under the Dii structure, pair up into two doubly degenerate modes under the latent space group D&.Therefore one would predict the b,, and b,, spectra, which correlate with the e irreducible representa- t Fig. 3 Crystal structure' of Cr(CO), projected down the Y-axis. The non-primitive translations parallel to Z and X axes necessary to bring the two hexagonal sub-structures to register are indicated for one molecule. Translations needed for the other molecules are obtained simply by transforming these translational vectors using the 0:: symmetry elements. tion of D&,, to be the same to a first approximation, and similarly, for the a, and b,, pair, which correlates with the e2, irreducible representation of @h. However, the molecular environment in the real structure has a contribution to its potential from a tetragonal distribution of molecules as indi- cated in Fig.l(c). The effect of this D,&hlatent-symmetry factor group is only to split the e,, irreducible representation of D6h but to maintain the degeneracy of el, (Table 1).There-fore b,, and b,, spectra are predicted to be much more similar to each other than the a, and b,, pair. This is exactly what is observed experimentally as reported by Adams and Taylor in ref. 8, Fig. 4 and 8. However their assignments of a band at 2009 cm- to the b,, factor group component of the v, ,molecular e, ,mode and that at 2024 cm- ',the strongest b,, symmetry band in this spectral region, to a combination (v6 + v,) disagree with the present analysis.Adams and Taylor' have expressed their surprise at this apparent inten- sity anomaly but appear to have stuck to their assignment on the assumption that v, should show only very little factor- group splitting. Therefore we prefer the earlier assignment of Kariuki and Kettle' of the most intense b3e band (ca. 2024 cm-') to v, . With this assignment, the factor-group com- ponents of v, in the Raman fall into two pairs, b,, and b2g very close in frequency (at 2009 and 2007 cm-', respectively) approximately maintaining the el, degeneracy of the latent space group while the symmetries a, and b,, are well separat- ed (at 2003 and 2024 cm-', respectively) showing the predict- ed destruction of the corresponding e2g degeneracy due to the distortions relating the real and latent symmetry structures.Use of Table 1 provides similar predictions for the IR spectra. The IR activity under & is confined to a2, and el, irreducible representations and the latter remains degenerate under the tetragonal symmetry and correlates with b,, and J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 b,, under the real factor group, Diz. Therefore b,, and b,, spectra are predicted to be similar to each other, but different to the b,, spectrum, in agreement with the reported IR spectra of Adams and Taylor in Fig. 2 of ref. 8. Distortion-induced IR and Raman intensities due to for- bidden modes of molecular units at polar site symmetries in crystals have been discussed by Jain and Bhattacharjee.' They have used perturbation theory and group theory to make predictions within the site group approximation.The assignment of a band at 1964 cm- ' in Cr(CO), (band h in ref. 5) to a genuine factor-group component derived from the molecular t,, mode (vg) by Kariuki and Kettle,' in accord with some earlier isotopic substitution work,,' is in agree- ment with Jain and Bhattacharjee's' , distortion-induced activity predictions. However, the molecular t,, mode is expected to give rise to six factor-group modes in the Raman, 2a, + b,, + 2b2, + b,, (Table 1). Jain and Bhattacharjee's', treatment does not provide an explanation as to which of these factor-group components would show the most inten- sity. The latent symmetry approach presented in this paper provides a criterion for such a distinction between the factor- group components.The correlation given in Table 1 shows the clear preservation of geradelungerade separation from the molecule to the hexagonal latent space group, and if this is strictly valid the molecularly IR-active modes would have no predicted Raman factor-group components and vice uersa. This is mainly what is observed experimentally, except for a very few exceptions. These exceptions however find an expla- nation in the intermolecular tetragonal latent symmetry space group. As already discussed, the irreducible representa- tions a, and b,, which correlate with e,, under D6h are the more sensitive modes to the intermolecular tetragonal poten- tial contribution, and are therefore predicted to be the most intense factor-group components of the molecularly for-bidden modes in the Raman.These predictions are in agree- ment with all experimental observations for this whole class of compounds.'*' Similar considerations would predict the distortion-induced IR activity of the molecularly only Raman-active modes to show intensity mainly in b,, sym-metry. Conclusion Here we provide an example of the latent symmetry approach to vibrational spectroscopy in the solid state which provides a physical realisation of the vibrational spectral con- sequences of intermolecular interactions present in the solid state, and thus an understanding of some relatively compli- cated spectral patterns of an important set of complexes, the hexacarbonyls of the Group 6 metals.This technique, which uses mainly symmetry-based arguments, provides an under- standing of these systems which is ostensibly independent of a particular mathematical model. This method clearly extends one's understanding beyond the oriented gas approximation' and is also shown to extend to the factor- group components of the distortion-induced IR and Raman intensity predictions due to forbidden modes of molecular units.', Further, explanation of the subtleties of the solid- state spectra in this way provides a clear picture of the important perturbations present, which would be of value in the determination of a meaningful mathematical modelling of these systems. References 1 (a) A.Whitaker and J. W. Jeffery, Acta Crystallogr., 1967, 23, 977; (b) B. Rees and A. Mitschler, J. Am. Chem. Soc., 1976, 98, 7918. 2 N. J. Hawkins, H. C. Mattraw, W. W. Sabol and D. R. Carpen-ter, J. Chem. Phys., 1955, 23,2422. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 1269 3 4 (a) F. A. Cotton and C. S. Kraihanzel, J. Am. Chem. SOC.,1962, 84, 4432; (b) L. H. Jones, R. S. McDowell and M. Goldblatt, Inorg. Chem., 1969,8,2349. D. M. Adams, W. S. Fernando and M. A. Hooper, J. Chem. SOC., Dalton Trans., 1973,2264. 10 11 12 U. A. Jayasooriya, S. F. A. Kettle and S. Mahasuverachai, J. Chem. Phys., 1987,86, 3127 and references therein. H. Poulet and J. P. Mathieu, Vibrational Spectra and Symmetry of Crystals, Gordon and Breach, Pans, 1976. E. B. Wilson, J. C. Decius and P. C. Cross, Molecular Vibrations, 5 6 7 8 D. A. Kariuki and S. F. A. Kettle, Inorg. Chem., 1978,17, 141. W. Scheuermann and K. Nakamoto, J. Raman Spectrosc., 1978, 7, 341. W. Rudorff and U. Hofmann, 2. Phys. Chem. B,1935,28,351. D. M. Adams and I. D. Taylor, J. Chem. SOC.,Faraday Trans. 2, 1982,78,1051. 13 The Theory of Infrared and Raman Vibrational Spectra, McGraw-Hill, London, 1955, ch. 3. Y. S. Jain and R. Bhattacharjee, J. Phys. C., Solid State Phys., 1985,18,5299. 9 A. Kastler and A. Rousset, J. Phys. Radiat., 1941, 2, 49; Vibra-tional Intensities in Infrared and Raman Spectroscopy, ed. W. B. Person and G. Zerbi, Elsevier, Amsterdam, 1982. Paper 3/07305J;Received 10th December, 1993.

ISSN:0956-5000    

DOI:10.1039/FT9949001265

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(9), 1271-1278 Valence and Core Photoemission of the Films formed Electrochemically on Nickel in Sulfuric Acid Yuanling Liang and Peter M. A. Sherwood" Department of Chemistry, Willard Hall, Kansas State University, Manhattan, KS 66506,USA Dilip K. Paul Department of Chemistry, St. Mary College, Leaven worth, KS 66048,USA The nature of the anodic films formed on nickel in 0.5 mol dm-3 sulfuric acid after polarization in the passive and transpassiveo state has been investigated by X-ray photoelectron spectroscopy. It is found that the film thickness (10-25 A) grows linearly with potential in the passive region. Film thickness drops at the beginning of the transpassive region, due to the further oxidation of NiO to Ni0,H (which may be soluble in the electrolyte).The results from anaerobic cell and ex situ experiments were compared in order to study the effect of atmo- spheric conditions on the electrochemical treatment. Experiments performed in the anaerobic cell show strong sulfate adsorption and the trapping of sulfate ions in the passive film. No significant sulfate adsorption is found in the ex situ experiments, presumably because the air-formed oxide layer prevents these ions from being trapped in the film. The passivation of nickel in sulfuric acid solution has been studied in detail for many years with different techniques.'-' ' Most of the published data show that the electrochemical behaviour of nickel depends upon the time of polarization and the pretreatment of the electrode.The passive layer formed is a mixture of, or a bilayer structure of, NiO and Ni(OH),. With an increase in potential the composition of the film changes with the NiO content in the film increasing with a surface analysis with XPS, the correlation of the electro- chemical behaviour with the nature of the species on the surface is achieved. To avoid artifacts such as oxide growth during exposure to the oxygen of the laboratory atmosphere, sample transfer from the electrolyte to the UHV system was performed in a specially designed anaerobic chamber. The use of ultra-high purity (UHP) argon as a protective gas pre- vents any further oxidation. On the other hand, a number of decline in the amount of Ni(OH), comp~nent.~*~-~~,~~-~ reactions occur under conditions where the electrode may The thickness of the layer is in the range of 1.1-1.7 nm deter- mined by coulometry,' electrochemical methods,' nuclear microanalysis and X-ray photoelectron spectroscopy (XPS).*-''*' Ho wever, there are still many unsolved ques- tions. Some controversies still exist among various investigators concerning the formation, composition, structure and proper- ties of the passive layer, despite several careful investiga- tions.In particular: (i) Nature and thickness of the passive film; some authors have reported a linear increase in film thickness in the passive region with p~tential.'*'~ Other authors suggest that the thickness of the passive layer was potential-independent.'*' (ii) The presence of sulfate ions in the passive film; we found sulfate ions on the electrode at almost all potentials, and a very small amount of sulfide present in the active and prepassive regions.Some previous workers have indicated that nickel sulfate exists as a constitu- ent of the passive Droste and Feller3 and Marcus and Grimalg have suggested that the sulfate was a kind of contamination on the passive film surface, the HS0,-orSob2-ions being adsorbed by the oxide film. In order to understand this system, surface analysis is needed in addition to electrochemical methods. Surface analysis methods in ultra-high vacuum (UHV), like XPS, compliment several in situ experiments in providing detailed chemical In XPS studies the core level spectra may become ambiguous due to the overlapping of a large amount of multiplet splitting and other satellite peaks.In such cases the valence band spectrum may become a powerful tool for the determination of the chemical nature of the surface layer. The aim of this work was to examine the passive layer in detail after a variety of electrochemical treatments, including treatment in the preactive, passive and transpassive regions. By combining the controlled electrochemical treatment and have either an adsorbed layer of oxygen or a very thin oxide film. The film may play a role in the electrochemical reactions either directly by modifying the energetics of adsorption of reactants or intermediates, or indirectly by changing the potential distribution at the surface.Thus a set of parallel experiments was also performed on the bench (the ex situ experiment) to investigate the effect of air exposure on the electrode surface chemistry. Experimental Instrumentation XP spectra were collected on a VSW HA10 spectrometer with a second UHV system used for electrochemical experi- ments. The design of this chamber and the operation of its \ I. 1.5 1.o 0.5 0.0 -0.2 potential/Vvs. SCE Fig. 1 Cyclic voltammogram of nickel metal in 0.5 mol dm-3 sul- furic acid solution with a scanning rate of 40 mV s-' 1272 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 1 Curve-fitting parameters for Ni 2p3,, and 0 1s core XPS regions ~ ~~ Ni 2p reference in addition to this main sat. 1" sat.2" sat. 3" 0 1s work Ni metal 852.4 858.8 19, 34, 37-39 1.3' 5.2' NiO 854.1 2.2' 855.6 2.5' 860.9 5.P 864.4 5.P 529.9 1.5' 19, 35-39 Ni(OH), 856.4 2.6' 858.0 2.7b 862.5 5.P 531.8 1.9' 19, 34, 36-38 H,O -- - 533.6 1.9' 14,38 NiSO, 856.3 2.4' 858.2 2.7' 861.3 4.9' 532.3 1.9' 37, 38 " Sat. = satellite peak. ' FWHM. electrochemical cell has been described elsewhere.2 1,22 Both UHV systems could achieve a base pressure of lo-'' Torr. All XPS data were recorded in the fixed analyser transmis- sion (FAT) mode with a pass energy of 25 eV for core regions and 50 eV for the valence band region. Achromic Al-Ka radi- ation (1486 eV with a linewidth of ca. 0.85 eV) was used at a power of 250 W (10 kV, 25 mA).Data were collected with at least 17 points per eV to be sure to identify any subtle fea- tures that might be lost at lower resolution and larger step size. The spectrometer energy scale was calibrated using copper.32 Peak positions were referenced to the C 1s peak (284.6 eV), due to residual hydrocarbon on the sample surface. Argon-ion sputtering was carried out using of a B21 saddle-field ion source sputter ion gun operated at 2.8 kV and 1 mA. Ultra-high-purity argon was used for etching and as an inert atmosphere in the anaerobic cell experiments described below. Two three-electrode cells were used for anaerobic cell experiments and ex situ experiments, respec- tively. Potentiostatic polarization and cyclic voltammetry A B C D 175 167 537 529 880 856 40 8 binding energy/eV Fig.2 XP spectra for nickel metal polarized to different potentials in sulfuric acid (0.5 mol dm-3) solution within the anaerobic cell. Spectra are shown for the S 2p, 0 1s and Ni 2p core regions and the valence band region. The curve-fitting results are shown for the 0 1s region. (a) Shows an etched nickel metal sample. (b)-(f)shows results of polarizing a nickel electrode to different potentials for 15 min. All potentials other than -0.6 V were prepolarized to -0.6 V for 5 min; (b) -0.6, (c) +0.7, (6)+ l.l,(e) +2.0, (f) +2.5 V. were performed using a Thompson Ministat Research Poten- tiostat (model 402R). Sample Preparation Nickel foil from Alfa (of 99.994% purity) was used as the source of the nickel electrode. 0.5 rnol dm-' sulfuric acid solution was prepared from the concentrated acid (with 95.0%-98.0% concentration) obtained from Fisher (of an ACS specified purity) and quadruply distilled water.All potentials were measured and quoted with respect to the saturated calomel reference electrode (SCE), which is 0.245 V with respect to standard reversible hydrogen electrode, in 0.5 rnol dm-3 sulfuric acid solution. All of the sulfuric acid solu- tions and washing water were deaerated with UHP nitrogen for at least 8 h and then exposed to UHP argon for 10-15 min. A B C D ' -., ., . . i.* '.<..<;;<*.:, ,,..,. . :' v 172 164 534 526 880 856 40 8 binding energy/eV Fig.3 XP spectra for nickel metal polarized to different potentials in sulfuric acid (0.5 mol dm-3) solution for ex situ experiments. Spectra are shown for the S 2p, 0 1s and Ni 2p core regions and the valence band region. The curve-fitting results are shown for the 0 1s region. (a) Shows an etched nickel metal sample. (b)-(f)show results of polarizing a nickel electrode to different potentials for 15 min. All potentials, other than -0.6 V, were polarized to -0.6 V for 5 min; (b) -0.6,(c) +0.7,(d) +l.l,(e) +2.0,(f) +2.5 V. J. CHEM.SOC. FARADAY TRANS., 1994, VOL. 90 A B h v)Y.-C 862 852 862 852 binding energy/eV Fig. 4 Curve-fitting results for the Ni 2p,,, core XPS region for nickel polarized to different potentials in sulfuric acid (0.5 mol dmP3) solution for 15 min in both anaerobic cell and ex situ experiments.All potentials, other than -0.6 V, were prepolarized to -0.6 V for 5 min. (a)-0.6,(b) +0.7,(c) + 1.1, (6)+2.0, (e) +2.5 V. The nickel metal samples were polished mechanically with alumina (44pm). No trace of aluminium was found on the polished metals following XPS examination. They were also degreased with acetone and cleaned with quadruply distilled water. The metal thus had only an air-formed film on the surface. Two types of experiments were conducted. The anaerobic cell experiments were performed in the special designed anaerobic chamber discussed previously.2 'st2 These experi- ments involved an additional cleaning procedure for the samples by argon-ion sputtering for 1.5-2 h on each side with the surface purity monitored by XPS.The sample was rotated during the argon-ion bombardment to allow even treatment of the sample. This further cleaning procedure enabled us to remove all of the contamination and oxide resi- dues formed during exposure to air, resulting in reliable and reproducible starting conditions. The clean sample was then exposed to a positive pressure of an UHP argon atmosphere in the anaerobic cell before the electrochemical treatment. For the ex situ experiments, the sample, cleaned as described above except for the argon-ion sputtering, was immediately immersed in the deaerated 0.5 mol dm-, sulfuric acid and polarized. In all of the electrochemical treatments, the samples were prepolarized at -0.6 V for 5 min to remove oxide film, then pulsed to a chosen potential: -0.6, +0.7, +1.1, +2.0 and +2.5 V,respectively, and maintained there for 15 min.All the experiments were carried out at room temperature. After fin- ishing the electrochemical treatment, the electrode was removed from the solution with the potentiostat still switched on, and then washed eight times with deaerated quadruply distilled water. Finally, the sample was dried by evacuation (anaerobic cell experiments), or by gently absorbing the water with adsorbent paper (ex situ experiments). Afterwards the sample was transferred to the XP spectrometer chamber. For ex situ experiments, the transfer procedure took <5 min. The anaerobic cell experiments eliminated exposure of the samples to any oxygen in either the gaseous or solution state, except for the oxygen produced from the electrochemical reaction itself.1273 Curve Fitting Table 1 gives all of the parameters used for curve fitting of the 0 1s and Ni 2p3,, spectra. The fitting of these spectra uses the approach that we have used previou~ly,~~.~~ and is consistent with the fitting of these regions by other workers cited in Table 1. The 0 1s regon was fitted to three or four peaks (vide infra) and the detailed curve fits of this region are shown in Fig. 2 and 3 (later). The Ni 2p3,, region had a non- linear background removed using the Tougaard method4' and the detailed curve fits of this region are given in Fig.4 (later). A non-linear least-squares curve-fitting program was used with a mixed Gaussian/Lorentzian (G/L) peak-The peak profile was kept the same in all cases. Exponential asymmetric tailing, to account for conduction band interaction (CBI) effects, was added to the basic G/L lineshape for nickel metal. All the other peak profiles remained symmetric. Ka,, X-ray satellites were also included in the curve-fitting routine. All of the fitting param- eters, such as peak position, peak width and peak separation between the main peak and the satellites, were chosen according to previous work and curve fitting results of stan- dard compounds (as indicated in Table 1). Results Electrochemistry Nickel exhibits a typical active-passive behaviour.The potentiodynamic polarization curves of nickel in 0.5 mol dm-3 sulfuric acid solution can be found el~ewhere.~*~-'~ Fig. 1 shows the cyclic voltammogram of nickel recorded in 0.5 mol dm-3 sulfuric acid solution with a scan rate of 40mV s-'. The polarization potentials used to treat the samples were chosen so that there was one in the cathodic region (-0.6 V), two in passive region ( +0.7 and + 1.1 V), and two in the transpassive region ( +2.0 and +2.5 V), respectively. Surface Analysis Anaerobic Cell Experiments Fig. 2 shows the typical XP spectra of a nickel electrode in 0.5 mol dm-, sulfuric acid solution as a function of electrode potential for treatment in the anaerobic cell. 0 1s spectra were curve-fitted with four peaks, which correspond to 02-, OH-, H20 and respectively.The OH- region also contains some water intensity.22 The fitting parameters are listed in Table 1, the fitting results are given in Table 2, and the approximate film composition is given in Table 3. Table 3 has been presented so that the reader can obtain a rough comparison between the information provided by the 0 1s and the N 2p regions. All the atomic ratio data assumes the very crude model of a uniform homogeneous surface layer, and are adjusted for differences in cross-section and analyser transmission function. The table explains how the percent- ages of the oxides are obtained from the percentage peak areas of Table 2. It should be noted that the ratio shown in column A is obtained assuming that NiO contributes one oxygen atom and Ni(OH), two oxygen atoms.The percent- ages of oxidized nickel species would not expected to be the same for the 0 1s region as the Ni 2p region since the Ni 2p electrons have a lower kinetic energy (635 eV) than the 0 1s electrons (956 eV) and are thus more surface sensitive. Since we know that the oxide lies in the outer region of the film we would expect the percentages of oxidized nickel to be greater J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 2 Curve-fitting results for the 0 1s and Ni 2p,,, core XPS regions 2” Ni metal NiO Ni(OH), 0,-OH-‘ H2O so,,-potential peak % peak % peak % peak % peak % peak % peak % ex situ -0.6 852.3 70.8 854.0 856.2 10.3 530.1 8.9 531.8 64.2 533.6 (0.0) (0.2) (0.5) (0.6) (0.3) (0.2) (0.02) (0.1) (0.0) (0.1) 1.3’ 2.2’ 2.5’ 1.5‘ 2.P 2.P +0.7 852.4 27.1 854.1 856.5 24.1 529.9 23.0 531.8 66.0 533.6 (0.0) (0.3) (0.2) (1.2) (0.3) (0.0) (0.0) (0.0) (0.0) (0.1) 1.3’ 2.2’ 2.6’ 1.5‘ 2.P 2.P +1.1 852.3 18.6 854.1 856.2 26.7 529.8 25.0 531.9 65.4 533.6 --(0.1) (0.3) (0.2) (1.3) (0.43) (0.0) (0.0) (0.0) (0.0) (0.1) 1.3‘ 2.2’ 2.6’ lSb 2.1‘ 2.1‘ +2.0 852.4 48.7 854.1 856.4 19.0 530.0 19.5 53 1.8 61.8 533.6 (0.0) (0.3) (0.2) (1.2) (0.3) (0.0) (0.0) (0.0) (0.0) (0.1) 1.3’ 2.2’ 2.6’ 1.5‘ 2.P 2.P +2.5 852.3 21.1 854.0 856.3 23.1 529.8 26.7 531.8 58.3 533.6 (0.0) (0.3) (0.2) (1*4) (0.4) (0.0) (0.0) (0.0) (0.0) (0.1) 1.3‘ 2.2’ 2.6’ 1.5’ 2.P 2.P anaerobic cell -0.6 852.5 83.1 854.0 856.3 --531.5 43.0 533.5 532.2 35.3 (0.0) (0.3) (2-4) (2.6) (3 (0.0) (-) (-1 (0.0)1.3’ 2.2’ 2.5’ 1.9’ 1.9‘ 1.9’ +0.7 854.4 45.5 854.0 856.4 530.0 3.0 531.7 34.7 533.6 532.3 48.2 (0.0) (0.2) (0.5) (1.4) (-1 0.0) (-1 0.0) (-) (-3 (0.2)1.3’ 2.2’ 2.6’ 1.5’ 1.9’ 1.9‘ 1.9’ + 1.1 852.4 27.4 854.1 856.5 530.0 10.6 531.7 37.9 533.7 532.3 37.9 (0.1) (0.6) (0.6) (1-7) (0.0) (0.0) (-1 (0.0) (-1 (-9 (0.0)1.3’ 2.2’ 2.6’ 1.5’ 1.9’ 1.9’ 1.9’ +2.0 852.4 59.6 854.1 856.3 529.9 2.1 531.6 39.2 533.6 532.3 43.4 (0.0) (0.3) (0.7) (0.8) (-3 (0.0) (-4 0.0) (-1 (-3 (0.0)1.3’ 2.2’ 2.5’ 1.5‘ 1.9‘ 1.9’ 1 .9’ +2.5 852.5 19.1 854.1 856.5 529.8 11.0 531.7 31.4 533.7 532.4 44.3 (0.0) (0.05) (0.4) (1.3) (-3 (0.0) (-4 (0.0) (-) (3 (0.0)1.3‘ 2.2’ 2.5’ 1.5‘ 1.9* 1 .9’ 1.9‘ a Peak = peak centre.% = % area of the Ni 2p or 0 1s core XPS region. The potential is in V us. SCE. ’FWHM. ‘This peak includes some H20intensity. in this region. However, the trends are the same for both also contains some intensity due to adsorbed water. Columns regions. The accuracy of the percentage of Ni(OH), is limited A and C should be roughly the same (as observed), and differ by the fact that its Ni 2p and 0 1s features are similar to by the amount of water in the film. The S : 0 atomic ratio is NiS04 and by the fact that the 0 1s region for hydroxide only accurate to about +5% due to the low intensity of the S Table 3 Approximate composition information from core XPS regions“ ~~ ~ ~~ 0 :Ni atomic S : 0 atomic 0 :Ni atomic 0 : Ni atomic % 0 1s as NiO (%) Ni(OH), (%) ratio from ratio from ratio ratio from sulfate Ni Q,,, S 2p/O 1s adjusted for 0 Is :Ni 2p,,, from S :0 Nimetal(%) Ni2p 0 1s Ni2p 0 1s curve peak ratio sulfate peak ratio atomic ratio potential 1 2 3 4 5 A B C D E ex situ -0.6 70.8 18.9 7.6 10.3 22.4 0.40 0.03 0.81 0.86 12+0.7 27.1 48.8 24.8 24.1 22.7 0.97 0.07 0.97 1.08 24 + 1.1 18.6 54.5 30.7 26.7 25.4 1.08 0.06 1.08 1.23 24 +2.0 48.7 32.3 16.6 19.0 21.2 0.70 0.03 0.79 0.85 12 +2.5 21.1 55.8 49.1 23.1 31.6 1.02 0.06 1.62 1.84 24 anaerobic cell -0.6 83.1 7.6 0 9.4 14.2 0.26 0.13 0.49 0.66 52 +0.7 45.5 24.3 4.9 30.3 28.1 0.85 0.14 1.17 1.62 56 + 1.1 27.4 42.6 18.0 30.0 32.2 1.03 0.12 1.29 1.7 48 +2.0 59.6 14.6 2.5 25.8 23.3 0.66 0.17 0.79 1.19 68+2.5 19.1 44.8 21.3 36.2 30.5 1.17 0.17 1.60 1.94 68 ~~~~ ~ “ Columns 1,2 and 4 are taken from the curve fit of the Ni 2p,,, region in Table 2. Column 3 is calculated by multiplying the percentage of the 0,-components in the 0 1s region of Table 2 by the 0 :Ni atomic ratio shown in column D.Column 5 is calculated by multiplying the percentage of the OH- component in the 0 1s region of Table 2 by the half the 0 :Ni atomic ratio shown in column D [since Ni(OH), has two oxygen atoms for each nickel atom]. Column 5 is further modified for the ex situ experiment by reducing the percentage of the OH- component in the 0 1s region of Table 2 by the percentage shown in column E since the ex situ data has no sulfate component fitted in the 0 1s region.Column C is obtained by adjusting the data in column D for sulfate not included in column A using column B. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 2p region. Changing the method from background subtrac- tion in the Ni 2p region will also change the percentages, the iterative non-linear meth~d~'.~~giving lower percentages than the Tougaard method4' used here, though the trends are the same. Cathodic Region. After argon-ion sputtering, the intensity of the 0 1s signal was extremely weak. No 0 2s signal is found in the valence band region (which is less surface sensi- tive due to the higher electron kinetic energy). This indicated that the surface is almost oxide free when the sample is exposed to the electrolyte. Cathodic polarization of the sample at -0.6 V gave some oxidized nickel.We believe that the oxidized nickel came from reaction with the deoxy-genated water, since the amount of oxidation was consistent with that found for clean metal samples exposed to deoxy- genated water in our previous studies2, in the anaerobic cell. Since the initial metal sample had no oxide present, and the electrochemical conditions would be expected to reduce any oxidized nickel species on the metal surface to metal, reaction with water would be expected to account for all the observed oxidation. Substantial amounts of adsorbed water and sulfate were indicated by the significant 0 1s peak and a weak 0 2s feature [due the lesser surface sensitivity of the valence band region (oide supra)].The S 2p region shows two peaks. The peak at higher binding energy (169.1 eV) corresponds to sulfate, and the peak at lower binding energy (163.5 eV) to sulfide. Sulfide results from the reduction of the adsorbed sulfate at the nickel surface. Passioe Region. In the passive region [Fig. 2(c), (41, the signal due to the oxidized nickel increases with increasing potential. For example, there is more NiO after polarization to 1.1 than to 0.7 V, which can be seen from the curve fitting results of the 0 1s and Ni 2p regions (Fig. 2B and Table 2). The percentage of Ni(OH), remains approximately constant as the polarizing potential increases. Ni(OH), could be slightly dissolved by sulfuric acid, and a significant amount of the Ni(OH), content may be formed after the sample was removed from the acid solution.After leaving the acid solu-tion all of the samples have the same environment, so it is reasonable to assume that they form almost the same amount of Ni(OH), . The increase in the intensity of the sulfate signal indicates that the amount of sulfate on the nickel surface increased with increasing potential (see Fig. 2A). At the same time the sulfide signal at lower binding energy is extremely weak and finally disappears after anodic polarization. Since the main peak of the Ni 2p signal from NiSO, and the first satellite peak (see Table 1) is at about the same position as that of Ni(OH), ,it is difficult to distinguish these compounds in the Ni 2p region. However, the NiSO, signal can be seen from the S 2p region and in the valence band spectra.The two small peaks at around 12.8 eV and 15.5 eV (Fig. 2D) are caused by sulfate. According to our standard spectra4, the separation between these peaks is about 2.7 eV for sulfate and around 3.9 eV for hydrogensulfate, and the relative inten- sities of the two peaks are different. For sulfate the peak at higher binding energy is more intense than the one at the lower binding energy, but the situation is reversed for hydro- gensulfate. The shoulder of the Ni 3d signal makes the rela- tive intensity of these two peaks less easy to determine, but the separation between the two peaks is clear and this allows us to identify the sulfur as sulfate and not hydrogensulfate.The 0 2s region is dominated by oxygen originating from sulfate and hydroxide. Transpassioe Region. In the transpassive region, the signal due to oxidized nickel decreases significantly at 2.0 V and then increases at 2.5 V. From both the 0 1s and Ni 2p regions (Table 2) we can see that the most significant change occurs for the NiO component, the change of Ni(OH), com-ponent being less than 7%. This result suggests the further oxidation of NiO and/or Ni(OH), to form NoO,H, which can be dissolved by sulfuric acid and would result in the thin- ning of the passive layer. The oxidation and reduction of the Ni(OH),/NiO,H pair has been discussed elsewhere as a main process for the nickel The small change in the relative intensity of the Ni(OH), component further supports the suggestion that most of the Ni(OH), was formed after the nickel electrode was removed from the acid solution.At 2.5 V oxygen was vigorously released and the formation of NiO becomes the dominant process again, resulting in the second- ary passive region. In this set of experiments, the sulfate signal can be seen in both the core level and valence band spectra. The washing procedure was the same as that used in our earlier work,,, where it was found that three washings removed all of the soluble sodium and chloride species after immersion of nickel into sodium chloride solution. In this work we washed each sample eight times. However, there was still a significant amount of sulfate left on the nickel surface as described above.It has been suggested that sulfate is a kind of contami- nation layer adsorbed on the oxidized nickel surface. Others have suggested that the sulfate is incorporated within the oxide layer, since it cannot be removed by washing the sample. Since NiSO, is a soluble compound, and our pre- vious work suggests that the cleaning procedure can remove soluble species on the nickel surface, it is unlikely that the sulfate exists as a physically adsorbed layer on the oxidized nickel surface. Ex situ Experiments Fig. 3 shows the XPS results for the ex situ experiments, with the curve-fitting data being shown in Table 2, and the approximate films composition in Table 3. The results show almost the same oxidation pattern as that of the anaerobic cell experiments (note the Ni 2p and 0 1s regions in both experiments).Cathodic Region. Cathodic reduction at -0.6 V shows some oxidation [NiO and Ni(OH),] on the electrode surface. As in the anaerobic cell case we suppose that this is caused by reaction with water, but one notes that the oxidation level is about twice as high, presimably because not all the air formed oxidation was removed by electrochemical reduction in this region. The fact that the NiO concentration is nearly twice as great as the Ni(OH), concentration supports this contention. Passive Region. In the passive region an increase in anodic polarization potential causes the intensity of the oxidized nickel signals to increase and the intensity of the metal signals to decrease.NiO is accumulated in the passive layer. No significant change in the amount of Ni(OH), was observed. Transpassioe Region. In the transpassive region, the inten- sity of the oxidized nickel signals decreased at 2.0 V and increased again at 2.5 V. The difference in these results from those of the anaerobic cell experiments is that only a very weak sulfur signal due to sulfate (S 2p) was observed. No significant sulfur contribution to the valence band spectra was seen. The two-peak feature at 12.8 eV and 15.5 eV was not observed. It is reasonable to assume that the washing procedure for these two sets of experiments should not cause this significant difference in sulfate concentration if the sulfate was simply adsorbed by the oxidized nickel film.Nevertheless, in order to check the effect of the washing procedure, an additional experiment was done. In this experiment, the nickel sample was cleaned in the normal manner (oide supra) and then argon-ion sputtered to remove all of the oxide layer. After this cleaning procedure the nickel sample was removed from the UHV chamber and electrochemically polarized imme- diately. This procedure was kept exactly the same as that of the anaerobic cell experiments, except that the nickel sample was exposed to air before electrochemical treatment. The sample was dried in air after treatment, and transferred into the UHV chamber within 5 min. This experiment gave the same result as that of the normal ex situ experiments, namely an extremely weak sulfate (S 2p) signal, and no sulfate signal in the valence band and other core regions. Details of the Curve Fitting of the Ni 2p,,, Region Fig.4 gives the curve-fitting results for the Ni 2p,,, region. All fitting parameters are given in Table 1, which uses peak positions and peak widths obtained from the spectra of stan-dard samples and the literature. The peak intensities were variable values determined by the fitting program. Three components, nickel metal, NiO, and Ni(OH),-NiS04 were used to fit all of the spectra. NiSO, cannot be distinguished from Ni(OH), ,since both of them have a Ni 2p,,, main peak and first satellite peak at almost the same binding energy. There is, however, a difference in position of the second satel- lite peak for these two compounds (Table 1).The fitting results are given in Table 2. We can see that the ex situ experiments always exhibit more oxidation than that of anaerobic cell experiments at the same potential, except at 2.5 V. We believe this to be caused by the additional oxida- tion of nickel by air in the ex situ case. Column B of Table 3 gives the S : 0 atomic ratio. This ratio is very small and has an almost constant value. This ratio is about three times greater in anaerobic cell experi- ments than the in ex situ experiments. Discussion Differences between the Anaerobic Cell and ex situ Experiments It is important to reconcile the differences between the ex situ and anaerobic cell experiments in the amount and ease of removal of surface sulfate.The experiments above suggest that the difference arises from the fact that the ex situ experi-ments have an air-formed oxide/hydroxide layer whereas the anaerobic cell experiments have an argon-ion etched metal surface. The argon-ion etching would be expected from previous studies of the oxidized nickel system48*49 to lead mainly to sputtering (for example 400 eV argon ions lead to about 70% sputtering and 30% reduction). Our saddle field gun provides uniform etching, though it is possible that some surface roughening occurs. Certainly our XPS studies show that the etched surface is almost completely metallic. We believe that in the anaerobic cell experiments sulfate ions were chemisorbed by the clean nickel surface.Sub- sequent electrochemical oxidation leads to the formation of an oxidized layer over the adsorbed sulfate layer. The sulfate ions were thus trapped between the nickel metal and the passive layer. This would explain why the sulfate ions cannot be washed away and thus they contribute significantly to the valence band and core level spectra. In the ex situ experi-ments, the sulfate ions are adsorbed onto an already air- oxidized metal. In this situation these ions are easily removed by washing. A second difference between the ex situ and the anaerobic cell experiments is that the former experiments show more adsorbed water and hydroxide ions in the valence band J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 region. This is evident from the more intense 0 2s signal over most of the potential range for the ex situ case (compare Fig. 2 and 3). Our work indicates that we cannot obtain a clean nickel metal surface by simple electrochemical reduction treatment, even at -0.6 V which is far below the redox potential of nickel metal. This reductive treatment still gives a significant amount of NiO and Ni(OH), in the case of the ex situ experi-ments, and a small amount of nickel hydroxide and/or sulfate in the anaerobic cell experiments. As discussed above this oxidation is almost entirely caused by reaction with water in the anaerobic cell case, combined with air oxidation during transfer in the ex situ experiment.Some residual strongly che- misorbed water may result in both cases. Estimating the Thickness of the Oxidation Layer It is possible to obtain an estimate of the oxide layer thick- ness using a simple intensity model, though one should make clear that such an approach assumes a uniform surface and a homogeneous mixture of NiO and Ni(OH), ,which we point out below is a rather crude approximation.50i51 This simple model uses an expression derived from the familiar intensity expressions where the intensity ratio of metal to oxidized nickel is given by where N, and N, are the number fraction of nickel atoms of NiO and Ni(OH), ,respectively. In this case we can assume a,,, = a,, (a,,, and a,, are the photoelectron cross-sections of the metal and the oxide, respectively), K the spectrometer factor, K, is obtained from the spectra using the previously described method48 as 1.36 for Ni(metal)/Ni(OH), and 0.98 for Ni(metal)/NiO. r~ is the reciprocal of the mean escape depth (A), D is the density of the atom in the material under investigation and x is the thickness of the film. Table 4 shows the values of G, D and molecular weight for this calcu- lation.8.1O Results of this evaluation for nickel samples polarized at different potentials are given in Fig.5. In the passive region the oxide thickness increases linearly with the electrode potential. The thickness fell around 2.0 V and increased again at 2.5 V. The oxide layer for the ex situ experiments was always thicker than the anaerobic cell thicknesses, except at 2.5 V.The minimum thickness amounts to 4 A arising from the film formed by adsorbed water and/or sulfate species. Previous angle-dependent XPS studies,, show (as expected) Table 4 Calculation parameters for the passive layer thickness &' Dlg cm -molecular weight /g mol-' Ni 0.18 8.90 58.71 NiO 0.10 6.67 74.71 Ni(0H) 0.073 4.15 92.7 1 NiOOH -4.68 91.71 J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 2.5 -2.0 -E5. 1.5-v) tY .-u 1.0-5 0.5 -0.0 4 I 1 I I I I i -1.0 -0.5 0.0 0.5 1.0 1.5 2.0 2.5 potentialp vs. SCE Fig. 5 Change in the total thickness of the oxidized nickel layer, formed in 0.5 mol dm-3 sulfuric acid solution, with the polarization potential.The polarization time is 15 min. that it is more likely that NiO is found in the inner part of the film, with an accumulation of hydroxide and water in the outer parts of the layer. The sulfate ions may exist either in the inner part (anaerobic cell experiments) or outer part (ex situ experiments) as discussed above, though this is dificult to determine by angle-dependent XPS. Clearly the film is much more complex than the simple homogeneous model that we have used to obtain these thickness values, and so such values should be regarded as a crude estimate of the film composition. Conclusions Comparison of the data from both anaerobic cell and ex situ experiments for nickel oxidation illustrates the importance of conducting such experiments under anaerobic conditions.Air-formed oxide films on the metal cannot be removed by electrochemical reduction, and this film can substantially affect the amount and ease of removal of sulfate ions. The use of this approach has allowed us to support and extend the conclusions of our previous study.8 The main chemical changes resulting from electrochemical oxidation are: (i) the film present on the electrode surface in the passive region is composed of NiO and Ni(OH),, the former being the passive The average thickness of this layer is in the region over the potential region where passivation occurs. On increasing the potential, the thickness of the passive layer increases linearly. (iii) At the beginning of the transpassive region, NiO and/or Ni(OH), could be further oxidized to form NiO,H, which can be dis- solved by acid solution and result in the breakdown of the passive layer.At higher potential (about 2.5 V) vigorous evolution of oxygen occurs, and a second passive region occurs. (iv) The adsorption of sulfate ions depends upon the pretreatment of the nickel electrode. Sulfate ions be chemi- cally adsorbed on a clean nickel metal surface, and then be subsequently trapped within the passive layer. Electrochemi- cal treatment of metal in the ex situ experiment where the metal has an air-formed oxide layer probably leads to adsorption of sulfate by the oxide layer, from which sulfate is easily removed by washing. This material is based upon work supported by the National Science Foundation under Grant No.CHE-8922538. The US Government has certain rights in this material. 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ISSN:0956-5000    

DOI:10.1039/FT9949001271

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(9), 1279-1284 Nickel incorporated into Anodic Porous Alumina formed on an Aluminium Wire Nobuyuki Ohji, Nobuhito Enomoto, Takanori Mizushima and Noriyoshi Kakuta Department of Materials Science, Toyohashi University of Technology, Tempaku, Toyohashi, Aichi 44 1, Japan Yoshio Morioka and Akifumi Ueno* Faculty of Engineering, Shizuoka University, Johoku, Hamamatsu, Shizuoka 432, Japan Porous alumina films have been prepared on the surface of an aluminium wire by anodic oxidation in an oxalic acid electrolyte. The pore size could be controlled by the anode potential, and the thickness of the alumina film was a function of the anodizing time as well as the anode potential used. Silica powders containing finely divided Ni metal particles were incorporated into the micropores of the alumina films, which were then used as wire catalysts.The wire catalyst did not require any supplementary heating during catalytic reactions, since after 1 min of applying a small voltage across the ends of the wire, a temperature of 573 K could be reached. The catalytic performance of the nickel particles mounted in the aluminium wires for the hydrogenation of but-l- ene was measured and compared with that of a conventional supported-nickel catalyst. Since the proposal of hexagonal pillar cells by Keller et al.' as a model for the micropores formed on anodic aluminium films, the morphology and mechanism of pore formation on anodic alumina fims have been of considerable interest in the fields of corrosion and surface modification science and tech- nology.Takahashi et d2and Furneaux et aL3 have devel- oped techniques for direct observation of these hexagonal pillar cells by transmission electron microscopy using an ultramicrotome for the preparation of the sample specimen, resulting in more precise measurements of the cell sizes and the pillar heights of the anodic films. On the basis of these measurements, the nucleation and growth of the porous anodic films on aluminium, prepared in various electrolytes, have been extensively discussed by Thompson et al.4-6 They reported that the local field strengths across the oxide (barrier) layers, varying with the electrolyte employed, control the steady-state anodizing behaviour observed in the acidic electrolytes under condi- tions of constant v01tage.~ In contrast, the application of these porous anodic films to novel functional materials has been of interest in the fields of magnetic data storage**' and catalysis," since these pores could work as nanotemplates for incorporation of small metal and/or metal oxide particles." In these studies, the method of incorporation of the small metal and/or metal oxide particles into the micropores formed on the anodic films is the key technique.In this work, anodic aluminas formed on an aluminium wire were immersed in sols consist- ing of ethyl silicate and nickel nitrate dissolved in ethylene glycol,'2 where nickel ions were trapped in the -Si-O-% Si-nets of the silica sols forming an -Si-O-Ni-O-% Si-structure.After reduction of the calcined aluminium wire in flowing hydrogen, some of the finely divided nickel par- ticles dispersed in the silica power were found to be incorpor- ated into the micropores of the anodic alumina films on the aluminium wire. Some kinetic studies on the hydrogenation of but-1-ene were carried out using these wire catalysts at various temperatures. It was found that the reaction required the same activation energy as over a conventional supported- nickel catalyst. Experimental Preparation of Anodic Porous Alumina Porous alumina films were prepared by anodizing an alu- minium plate (40 mm x 30 mm x 0.23 mm) or a spiral alu- minium wire (1 m x 0.25 mm od), both >99.99% purity obtained from Furukawa Electric Industry Co., at a voltage in the range 10-80 V in a 0.16 mol dmP3 oxalic electrolyte at 293 K using a graphite cathode and a regulator dc supply (Kikusui Denshi Co., PAD-1 10-5).The current observed was rapidly depressed in the fist 10 s of anodizing, and then gradually increased to a constant value (165 mA for a plate and 51 mA for a wire at 50 V, respectively) within 100 s. Unless otherwise specified, anodizing was continued for 10 min at constant voltage. The porous alumina films formed on the aluminium plates were dried and calcined at 623 K for 1 h, and then placed in a high-resolution scanning electron microscope (SEM, Hitachi S-900), operated at an accelerating voltage of 5 kV, in order to measure the open pore sizes and the film thicknesses (pillar heights), respectively.Since SEM observations of the porous films on the aluminium wires was difficult, the pore sizes and film thicknesses on the spiral wires were postulated to be the same as those observed on the plates under the same anodic conditions. Incorporationof Nickeldispersed Silica into the Micropores The anodic porous aluminas formed on the plates and/or wires, dried and calcined at 623 K in air for 1 h, were immersed in sols, prepared by hydrolysis of a mixture of ethyl silicate (30 an3)and nickel nitrate (8.9 g) dissolved in ethyl- ene glycol (150 an3)with a small amount of nitric acid (2 cm3).12 The sols consisted of silica colloids, trapping nickel ions in the -Si-0-Si- nets forming an -Si-0-Ni- 0-Si-structure evidenced by extended X-ray absorption fine structure (EXAFS) measurements.' The silica colloidal particles were roughly evaluated to be <10 nm in diameter on average.14 After immersion for 10 min at 298 K, the plates and/or wires were withdrawn at a pulling rate of 80 mm min-', followed by drying and calcining at 623 K for 4 h, and by reduction at 673 K for 1 h in flowing hydrogen.The formation of nickel metal particles in the porous alumina on an aluminium plate was posed by X-ray thin-film diffraction (Rigakudenki, RAD-SA), with an incident angle of 0.5 and a graphite slit system in front of a detector. A Cu tube was operated at 50 kV and 40 mA. In the present work, silica- coated wire catalysts were also prepared by immersing cal- cined anodic wires into a silica-sol solution consisting of ethyl silicate and a small amount of nitric acid.The anodic alumina films formed on a plate were mounted in a resin and sliced by an ultra-microtome into thin films, which were immersed in an aqueous solution of mercury(r1) chloride.” The sliced alumina films were separated from the aluminium metal plate owing to the formation of an Hg-A1 amalgam during the immersion. The alumina films separated from the aluminium plate were examined by tranmission elec- tron microscopy (TEM, Hitachi H-800 instrument) and elec- tron probe microanalysis (EPMA, Horiba EDX Instrument, equipped with an energy-dispersive detector), in order to measure the shapes and sizes of the nickel metal particles and to estimate the amount of nickel metal loaded in the porous alumina films, respectively.Catalytic Performance of tbe Wire Catalysts Although the alumina films formed on aluminium wire were readily heated to 573 K in 1 min simply by loading a small voltage to the ends of the wire, the films formed on the plate were heated to at most 323 K, because of the low electric resistance between both sides of the plate. In order to measure the temperature of the alumina films, a thermo-couple was attached to the wire surface using a heat-resistant resin, although temperatures >573 K could not be measured precisely. Hydrogenation of but-1-ene to butane was carried out at 358, 388 and 419 K using a closed-circulation system, equipped with a cylindrical glass reactor possessing two copper electrodes to which each end of the wire catalyst was connected to be heated electrically. Besides butane, butene isomers such as (E)-and (Z)-but-2-ene were detected during the reaction by gas chromatography using a column packed with VZ-7. The initial pressures of hydrogen and but-1-ene were both 53 kPa, and the rate constants for the formation of butane were estimated assuming the following rate equation : d(Pbutane)/dt = k(Pbut-l-ene)0.S(P~2)0’5.16For comparison, supported-Ni catalysts were prepared by immersing alumina powders in sols formed from hydrolysis of a mixture of ethyl silicate and Ni nitrate dissolved in ethyl- ene glycol, whose compositions were described above.The catalyst was then dried, calcined and reduced in flowing hydrogen in the same manner as mentioned above. The loading of Ni metal in the catalyst prepared was 2.0 wt.%, measured by X-ray fluorescence spectroscopy after extraction of the Ni ions with hot nitric acid, whereas the amount of Ni J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 ions in the alumina films, separated from the aluminium plates, was roughly estimated to be 1.9 wt.% by EPMA mea- surements. The catalytic performance for the hydrogenation of but-1-ene was also studied in a similar manner to that applied to the wire catalysts. Results Characterization of Anodic Alumina formed on Aluminium Plates The pore size, alumina film thickness and barrier layer of the anodic alumina films are defined in Fig.1, where the Keller model for anodized porous alumina films is drawn. High- resolution SEM photographs of the porous anodic alumina formed on an aluminium plate held at 50 V for 10 min, fol-lowed by calcination at 623 K for 1 h in air, are given in Fig. 2(a) (plane view) and (b) (sectional view), where the open-pore sizes and the film thickness (pillar heights) were measured. In Fig. 3 are shown the changes in average pore size and film barrier layer Fig. 1 Keller model for the anodic porous alumina films, showing the definitions of pore size, film thickness and barrier layer (a1 300nm (b) 300nm Fig. 2 SEM photographs of anodic alumina films: (a) plane view and (b) sectional view of the alumina films formed on the aluminium plate J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 400 -(a) C 0 0300 -0 0 20 40 60 80 voltagep 15 0 A E,10 --. IG 8 Y .-5 Ez5 8 0 20 40 60 80 voltage/V Fig. 3 Relationships between the voltage applied to the aluminium plate and (a) the pore size and (b) the alumina film thickness: (0) experimentally and (A) theoretically observed thickness on the aluminium plates anodized at various volt- ages for 10 min. Film thicknesses calculated from the follow- ing equation are also given in Fig. 3:” h=-Mlt 6Fd where h is the film thickness, M is the molecular weight of alumina, I is the current observed during steady-state anod- izing, t is the anodizing time, F is the Faraday constant and d is the current density at the aluminium surface.The apparent surface area of the aluminium plate used in this work is 24.3 cm2 (see Table 1) and the anodizing time is 10 min. The current during steady-state anodizing varied with the anode potential used (see Fig. 4). From the equation above, the film thickness should be proportional to the anodizing time under conditions of constant current. The results obtained at a con- stant voltage (50 V), and hence at a constant current (165 mA), are given in Fig. 5. In Table 1 the apparent surface area, the current during steady-state anodizing under 50 V, and the current density at the surface of the aluminium plate are Table 1 Apparent surface areas, currents during steady-state anod- izing at 50 V, and the current densities on the aluminium plate and wire apparent surface area current value current density /cm /mA /mA A1 plate A1 wire 24.32 7.62 165.0 51.0 6.78 6.70 compared with the corresponding values for the aluminium wire.Ni Metal Particles in the Porous Alumina Films of Aluminium Plates High-resolution SEM photographs of the reduced porous alumina on an aluminium plate are shown in Fig. 6(a) (plane) 1000 750 C a Ez2 500 0 L 30 250 0 0 0 0 001 I I I 1 I 0 20 voltage/\/ 60 8040 Fig. 4 Variation in current during steady-stake anodizing with applied voltage on the aluminium plate 15 0 0 0 0 0 0 0 20 40 60 anodizing time/mi n Fig.5 Linear relationship between alumina film thickness and the anodizing time J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 H 300 nmJb’ Fig. 6 SEM photographs of the hydrogen-reduced alumina films on the aluminium plate. Silica-supported nickel particles were dispersed onto the alumina films: (a)plane view and (b) sectional view. and (b) (sectional view), where the alumina surface was com- pletely covered with a nickel-containing silica film. It was also noted from Fig. qb) that small-grain silica powders, probably including nickel metal particles, were deposited on t-i 80 nm Fig. 7 TEM photograph of nickel or nickel oxide particles dis- persed inside the pores of alumina films formed on the aluminium 0 the pore walls of the alumina films.Fig. 7 shows a TEM photograph of the reduced alumina films separated from the aluminium plate, where the formation of nickel metal par- ticles, or nickel oxide crystallites, in the micropores of alumina films is evidenced. X-Ray thin-film diffractions of both the calcined and reduced alumina films are given in Fig. Table 2 Conversion and selectivity observed on the wire catalyst for the hydrogenation of but-1-ene at 419 K for 24 h selectivity (%) conversion catalyst (%) butane (Z)-but-2-ene (E)-but-2-ene wire catalyst 27.9 31.9 40.0 22.1 wire catalyst 0.07 trace trace trace without Ni Reaction carried out at 419 K. I 30 40 50 60 70 I I 1 I I29ldegrees Fig.8 X-Ray thin-film diffractions of the calcined (a) and reduced 2.4 2.5 2.6 2.7 2.8 (b) alumina films formed on the aluminium plate. Silica-supported 103 KIT Al,nickel and nickel oxide particles were in the alumina films. (0) Fig. 9 Arrhenius plot of the rate constants for the hydrogenation of Ni.(A) NiO, (0) but-1-ene on the wire catalyst; Ea = 2.3 kcal mol-’ J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 8, which suggests that nickel metal particles and/or the oxides are deposited on the surface of the alumina films as well. Hydrogenation of But-1-ene over Ni-dispersed Wire Catalysts Each end of the wire catalyst (1 m) was connected to the copper electrode mounted in the glass reactor. Prior to the introduction of the H,-but-1-ene gas mixture into the reactor, the catalyst was re-reduced by passing a current of 2 A through it under an H, atmosphere (corresponding to heating the catalyst to 523 K), followed by evacuation at the same temperature.The hydrogenation of but-1-ene on the wire catalyst was carried out at 419, 388 and 358 K, which corresponds to passing currents of 1.5, 1.3 and 1.1 A, respec- tively, through the wire catalyst, for 24 h. Gas samples were taken every 3 h during the reaction. Table 2 lists the results obtained after reaction for 24 h both on this wire catalyst and on the silica-dispersed (without Ni metal) wire catalyst. No reactions were observed on the silica-dispersed wire catalyst, suggesting that hydrogenation, as well as isomerization, of but-1-ene occurred on the Ni metal particles dispersed on the wire catalyst.The activation energy was 2.3 kcal mol-’, cal- culated from the results shown in Fig. 9. Discussion Control of Pore Size and Film Thickness of the Anodic Alumina Films The rapid decrease in current through the aluminium plate during the early stages of anodizing was due to the formation of alumina thin films on the plate surface, and the successive gradual increase in current was caused by the local disso- lution of the alumina film into the solution of oxalic acid electrolyte. Under conditions of steady-state anodizing, equi- librium between the formation and dissolution of the alumina films on the aluminium plate was achieved.” The pore size, which was rather smaller at the open pore [see Fig. 2(b)],was controlled only by the voltage applied between the alu- minium plates and the graphite electrode, large pores with a large voltage and small pores with a small voltage, as depicted in Fig.3. Although the nucleation mechanism has not yet been well defined, the pores were homogeneously gen- erated over the plate surface, as can be seen in Fig. 2(a).The anodic film thickness was also a function of the applied voltage, and increased exponentially as the applied voltage increased, which was well reproduced by the theoretical curve obtained from the equation given above. A linear relationship between the oxide film thickness and the anodizing time at constant voltage and current (50 V and 165 mA) was observed (see Fig.5), as predicted from the equation. Con- sidering the fact that the pore size was constant at any part of the pore, except at an open pore [see Fig. 2(b)],this suggests that the dissolution of alumina predominantly occurred at the bottom of the pores (barrier layer) during the steady-state anodizing, and hence, the growth rate of the anodic alumina films might be a function of the local field strength across the barrier layer.6 As can be seen in Fig. 4, the current passing through the aluminium plate during steady-state anodizing also increased exponentially as the applied voltage increased, which is almost the same as the relationship between the oxide film thickness and applied voltage (see Fig.3). The major factor controlling film thickness is, however, not the current itself but the current density (4, as expressed in the equation. The current densities were calculated from the currents observed during steady-state anodizing (see Fig. 4) and the apparent surface area of aluminium electrode employed (see Table 1). Therefore, the current density through the aluminium plate was almost the same as that through the aluminium wire used here, suggesting that the relationship between the alumina film thickness and the applied voltage on the alu- minium plate, given in Fig. 3, might be the same as the corre- sponding relationship for the aluminium wire. The pore size, however, is controlled only by the applied voltage for both the plate and the wire.Characterization of Ni Metal Particles dispersed in the Anodic Alumina Films Because of the repulsive forces caused by the electric double layer around alumina, metal cations cannot approach the alumina surface in acidic solution,18 which makes it difficult to deposit metal or metal oxide particles on the alumina surface, especially on the pore walls of the anodic alumina films. In most cases, electrodeposition was carried out using an aqueous solution of the metal salts of interest, buffered with boric acid.’.’’ In the present work, the calcined anodic aluminium plate and wire were immersed in silica sols containing Ni ions in their net frameworks as -Si-0-Ni-0-Si-structures, which might not be affected by the repulsive forces around alumina, since these macromolecules are neutral.14 It was further confirmed that Ni ions were rel- eased from the sol nets and migrated into the pores of the alumina films when these macromolecules were adsorbed by the alumina surface, leaving silica layers on the surface of the alumina films.I4 As can be seen in Fig.8, nickel oxide crystallites were actually formed in the alumina films on the aluminium plate and some of the oxide crystallites were reduced by hydrogen to nickel metal particles. Because of the low melting point of aluminium, the reduction was carried out at 673 K, which is not high enough to reduce all of the nickel oxide crys- tallites to metal. The open pores observed on the surface of the calcined alumina films [see Fig.2(a)] were completely covered with a layer of silica, probably containing nickel and nickel oxide crystallites [see Fig. qa)], although a lot of cracks and holes were noted on the silica layer. Silica deposi- tion was also confirmed inside the pores, as shown in Fig. qb). Since it was difficult to distinguish nickel particles from the silica deposits in these SEM photographs, the alumina films were separated from the aluminium plate and subjected to TEM observation. A few small black spots, not so clear, can be seen in the TEM photograph, shown in Fig. 7. These spots are probably nickel or nickel oxide crystallites, about 50 A in diameter, deposited inside the pores. Judging from the number of particles inside the pores, most of the metal and metal oxide particles were, however, in the silica layer covering the alumina surface, which is different from the results obtained in our previous ~0rk.l~ This difference can be ascribed partly to the differences in the amounts and strengths of acid sites on the alumina surfaces.Many strong acid sites were present on the surface of the alumina (spheres), assisting the Ni ions to leave the net framework of the silica sol. Catalytic Performance for the Hydrogenation of But-lene Although most of the metal and metal oxide particles were found in the silica layer formed on the alumina surface, the aluminium wire containing nickel metal particles was used as a catalyst for the hydrogenation of but-1-ene at 419, 388 and 358 K, respectively. These temperatures were achieved by passing current of 1.5, 1.3 and 1.1 A, respectively, through the 1284 J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 wire catalyst. As can be seen in Table 2, hydrogenation to butane and isomerization to (E)-and (Z)-but-2-ene took place over the wire catalyst, but not over the wire without nickel particles. It was also of interest to see if the current passed could have any effects upon the catalytic performance 2 3 4 H. Takahashi, M. Nagayama, H. Akahori and A. Kitahara, J. Electron. Microsc., 1973,22, 149. R. C. Furneaux, G. E. Thompson and G. C. Wood, Corros. Sci., 1978, 18, 853. G. E. Thompson, R. C. Furneaux and G. C. Wood, Trans. Znst. Metal Finish., 1977,55, 117. of the nickel, since the nickel metal particles were located so close to the aluminium wire that the local electric fields across the alumina films were expected to play some role in the catalytic properties of the metal.The activation energy for the hydrogenation of but-l-ene to butane on the wire 5 6 7 G. E. Thompson, R. C. Furneaux and G. C. Wood, Trans. Znst. Metal Finish, 1978,56, 159. G. E. Thompson, R.C. Furneaux, G. C. Wood, J. A. Richardson and J. S. Goode, Nature (London), 1978,272,43. G. E. Thompson and G. C. Wood, Nature (London), 1981, 290, 230. catalyst was estimated to be 2.1 kcal mol-' (see Fig. 9), which 8 D. AIMawlawi, and M. Moskovits, J. Appl. Phys., 1991, 70, is very close to the values of 2.0 and 2.5 kcal mol-', reported by Dibeler and Taylor" and by Twigg,I6 respectively.For comparison, alumina-supported nickel catalysts, prepared in the present work, were applied to the hydrogenation of but-l- ene in the same temperature range, resulting in an activation 9 10 11 4421. C. K. Preston and M. Moskovits, J. Phys. Chem., 1993,97,8495. N. Nourbakhsh, B. J. Smith, 1. A. Webster, J. Wei and T. T. Tsotsis, J. Catal., 1991, 127, 178. D. G. Goad and M. J. Moskovits, J. Am. Chem. SOC., 1989, 111, 9250. energy of 2.3 kcal mol-'. These results indicated that there 12 A. Ueno, H. Suzuki and Y. Kotera, J. Chem. SOC., Faraday are no effects of local electric fields on the catalytic per- formance of nickel metal particles dispersed in the wire cata- lyst, although further details, such as the turnover frequency on a nickel atom, have not been evaluated in the present work. 13 14 15 Trans. 1, 1983,79, 127. K. Tohji, S. Tanabe, Y. Udagawa and A. Ueno, J. Am. Chem. SOC.,1984,106,612. T. Fujiyama, M. Ohtsuka, H. Tsuiki and A. Ueno, J. Catal., 1987,104,323. S. Wernick, J. Electrodep. Tech. SOC.,1934,9, 153. This work was supported by Grants-in-Aid for scientific research from the Ministry of Education, Science and Culture 16 17 G. H. Twigg, Proc. R. SOC., London, Ser. A, 1941,178, 106. M. Nagayama and H. Takahashi, Nippon Kagaku-Kaishi, 1972, 850. of Japan. 18 19 H. R. Kruyt, Colloid Science, Elsevier, Amsterdam, 1948, vol. 1. T. J. Marks, Science, 1985,227, 881. References 20 V. H. Dibeler and T. H. Taylor, J. Phys. Chem., 1951,55, 1036. 1 F. Keller, M. S. Hunter and D. L. Robinson, J. Electrochem. SOC.,1953,100,411. Paper 3/074741; Received 20th December, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949001279

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(9), 1285-1291 Modification of the Electronic Structure of Pd by U Films: Chemisorption of COT Thomas H.Goudert and Carlos A. Colmenares University of California , Lawrence Livermore National Laboratory, Livermore, CA 94550,USA X-Ray and ultraviolet photoelectron spectroscopies (XPS and UPS), Auger electron spectroscopy (AES) and thermal desorption spectroscopy (TDS) have been used to study the coadsorption of CO and U on polycrystalline Pd. We investigated the modification of the electronic structure of Pd at low U coverages and the mode of CO chemisorption (dissociative vs. associative) as a function of surface U concentration. U adsorption results in the narrowing of the Pd4d band. At low U coverages the density of states (DOS) at the Fermi level decreases and the U 5f electrons are localized. The CO saturation coverage at room temperature decreases with increasing U surface concentration, while CO adsorption at -165 "C is less affected. This is attributed mainly to a decrease of CO chemisorption energy but also to blocking of Pd adsorption sites by U.A decrease in the heat of chemisorp- tion for CO is explained by a change of the electronic structure of Pd by U. At low dosage U itself loses most of its reactivity, probably because of U-Pd solid-state bonding and below a critical U surface concentration CO dissociation becomes an activated process. High-temperature reaction between CO and Ucovered surfaces leads to the partial transformation of U into a surface oxycarbide.In the presence of this compound some CO is chemisorbed strongly on the surface and is stable even at 300°C,in contrast to CO on Pd metal. A similar effect has also been observed for classical promoters such as the alkali metals. One goal of fundamental catalysis research is to relate the that above monolayer coverage U overlayers interact catalytic properties of surfaces to their basic chemical proper- strongly with the Pd substrate and that surface alloying takes ties and ultimately their electronic structure. Surface reacti- place at relatively low temperatures ( <200 0C).9Therefore, it vity of catalysts is often discussed in terms of geometrical' is conceivable that the adsorption properties of Pd surface and electronic factors.* The latter relate catalytic properties atoms are also changed.to measurable electronic quantities (bandwidth, DOS at the Fermi level, orbital symmetry et~.).~In this context it has been argued that catalytic properties of the late transition Experimenta1 metals (TM) are due to an interplay of highly correlated d electrons and s-p electrons, which are found in broad bands. The measurements were performed using a double-pass cylin- This theory was introduced as the 'electron interplay drical mirror analyser (CMA). UPS measurements were made model'.4 Actinides may be interesting systems to test this using He I1 (40.81 eV) excitation radiation produced by a model because some of them have highly correlated f elec-windowless UV rare-gas discharge source.The total trons (U, Np, Pu) in metallic systems while others do not (Th, resolution in UPS was 0.2 eV. XP spectra were taken using Am).' Unfortunately, despite the similarities between actinide Mg-Ka (1253.6 eV) radiation with an approximate resolution 5f electrons and late TM d electrons the chemical properties of 1 eV. Auger spectra were measured by a lock-in technique of the elements are very different: actinides do not behave using 3 keV electrons. like late TMs but like early TMs owing to their descent from Thin layers of U have been prepared in situ by magnetron actinium which is a refractory metal. In earlier work we sputter deposition. This method was preferred to the conven- found the high reactivity of U manifested by the exclusive tional evaporation technique because of the difficulties associ- dissociative adsorption of CO on U metal even at -200 oC.6 ated with the latter: U has a low vapour pressure at the However we also found that U loses part of its reactivity melting point and has a strong tendency to alloy with most when diluted in an alloy such as UNi,.6 The solid-state filament or crucible materials.A UHV compatible, water- bonding between U and Ni seemed to decrease the affinity of cooled magnetron sputter head was operated at a target surface U for CO. Such decrease of surface reactivity has also current of 5 mA, a bias voltage of 300 V and a gas pressure of been observed for early transition metals (Zr)7 and the rare- 0.5-1.5 Pa. As sputter gas we used ultrahigh-purity Ar earth metals.' It provides the necessary condition for study- (99.9999%),which was further cleaned over hot Ca at 130°C.ing the catalytic properties of initially reactive elements. Deposition times varied from 1 to 10 s. The target and sub- In this paper we will discuss whether further dilution of U strate were kept at room temperature. A shield was installed in a TM matrix will suppress its reactivity sufficiently to to expose only the sample to the Ar-U plasma and keep U make it part of a low-reactivity catalyst. We studied the CO contamination of the chamber as low as possible. The target adsorption on a polycrystalline Pd surface doped with U. In was a high-purity U disc (30 mm radius and 3 mm thick), such a system we have to consider two independent effects: which was cleaned before introduction into the vacuum the reactivity of U surface atoms themselves and the changed system by mechanical polishing and nitric acid etching. The reactivity of Pd surface atoms.We showed in an earlier paper purity of uranium overlayers was checked by XPS and AES. Oxygen was tO.l atom% as determined by AES, while carbon was not detectable by either technique. t This work was performed under the auspices of the US Depart-The substre.te was an ultrahigh-purity, polycrystalline Pdment of Energy by Lawrence Livermore National Laboratory under foil. It was cleaned by sputtering (1 h at 1 keV, 5 x Pacontract No. W-7405-Eng-48. $ Present address : Commission of the European Communities, Ar and 10 mA current) at 600 K and flashing to 1100 K. This European Institute for Transuranium Elements, Postfach 2340, procedure was repeated several times until no impurities were D-7500 Karlsruhe, Germany. detected on the Pd surface by AES and UPS.1286 Results Clean U Overlayers on Pd We first determined how U deposits on the Pd surface in the coverage range below one monolayer (ML). In a previous study we found that at the multilayer coverage U accumu-lates on the surface, partially diffuses into the near-surface region and forms a near-surface alloy with Pd.9 Fig. 1 com-pares the U 4f spectra for several U coverages. The surface coverage was determined by the ratio of the U,,, :PdMNN Auger emissions. Interpretation of these data, however, depends on the way U deposits on the surface. If U sponta- neously diffuses into the near-surface region to form a near- surface alloy the dosage should be described in terms of U concentration. If U stays at the top surface forming an over- layer, U dosage should be described by partial coverage and overlayer thickness.The U 4f data (Fig. 1)indicate that at the lowest coverage (U : PdAEs= 0.23) U indeed stays on the top surface: the inelastic background is strongly suppressed as seen by comparing the intensity at the high binding energy (Eb) side with the intensity at the low E, side, drawn in Fig. 1 as a thin horizontal line below the high E, side. This points to the absence of inelastic scattering of the photoelectrons before escaping from the surface.Because in this paper we are mainly discussing systems with low U content we will describe the U dosage in terms of fractional coverage. We estimate it, using Seah's formula" and the low coverage data of Fig. 1 with an AES ratio of 0.23, to correspond to ca. 0.4 ML coverage. We compared these data to the XPS- derived surface coverage, determined by the ratio of the U 4f7,, : Pd 3d,,, areas,l0," and obtained also a surface coverage of about 0.4 ML. However, precise determination of the surface coverage is impeded by the uncertainty of the information depth, which for the UoVv transition lies between 0.6 and 4 ML. Therefore, we will not discuss data in terms of a precise coverage but will rather follow their evolution with coverage.Amore quantitative analysis, involving ISS data, will be reserved for a future publication. Further information on the surface U is obtained from linewidth, shape and binding energy of the U 4f lines (Fig. 1). With increasing U coverage the binding energy decreases from a value found for U 4f,,, in UPd, (390 eV)" to 388.5 eV, which is typical for U metal. We conclude that at low coverage, U-Pd interactions predominate with U atoms dis- persed on the Pd surface, while at high coverage U forms a thick overlayer of U metal. Further indications for atomic dispersion at low coverage are obtained from the linewidth of )I#, I,,, I,,, I,,, I,,, ,,,I ,,,, ,,,, r c.-u)C C .-J. CHEM. SOC. FARADAY TRANS., 1994, VOL.90 the U 4f. It is narrow at very low and very high coverage showing that in both cases U is found in one well defined chemical environment with U-Pd and U-U bonding, respectively. Agglomeration of U in clusters or islands would have produced non-equivalent U atoms with slightly shifted E, . This actually happens at intermediate coverage (for U: Pd AES ratio of 1.05 and 1.92) and results in the broadening of the U 4f emission. The fairly symmetric shape at low coverage indicates a low local DOS at the Fermi level (EF)for the U atoms as discussed by Doniach and Sunjic,13 which we attributed previously to the localization of the 5f electrons. With increasing U coverage the U 4f emission becomes asymmetric pointing to the delocalization of the 5f electrons at higher U concentrations.Fig. 2 shows valence band spectra with increasing U dosages. The 4d valence band of pure Pd has a maximum value at the E, . At low coverages, U deposition results in the decrease of this maximum. The Pd band narrows, the Pd 4d intensity close to E, is suppressed and the d-band centroid is shifted to higher E,. Such findings can be understood in terms of filling of the Pd 4d band and of diluting Pd in a surface U matrix. The driving force for this is provided by the filling of the Pd 4d band in Pd metal.I4 At high U coverage the intensity at the EF increases again. The emission at the E, is now mainly due to the 5f states of the electropositive U which, as in U metal and most of its alloys, are delocalized. At very high U dosage the Pd 4d peak becomes narrow and symmetrical, and the U 5f emission at the EF further increases.At this multilayer stage the surface consists of a concentrated U phase with Pd impurities atomically dis- persed in it.9 UPS Study of CO Adsorption at Room Temperature Fig. 3 shows UPS He I1 spectra of U/Pd surfaces on which 10 Lt CO was adsorbed at room temperature. Fig. 3 (a)-@) correspond to Fig. 2 (a)-@). On pure Pd CO is only adsorbed molecularly, as shown by the lines at ca. 7-7.5 and 11 eV, which are attributed to the 17c/5a and 4a molecular orbitals. CO is not chemisorbed dissociatively on Pd." With increas- ing U concentration the CO signal decreases showing that less CO is chemisorbed on the surface.This may be due to three factors. First, part of the surface may be covered by U, '1'1' 1 ' 1 I Pd4d U 5f ..A Pd4d:la) I.l.fIl I 10 8 6 4 2 0 E,IeV UPS He I1 spectra of U overlayers on Pd. U: Pd(AES) =Fig. 2 (a)l~~,~l~~~~l,,,,l,~,,l,,,,l,,',i,,,,0.00,(b)0.20, (c)0.30, (d)0.50, (e) 1.86 and (f)3.58.U coverage: (a)420 410 400 390 380 370 360 0.00,(b)0.37, (c) 0.56 and (d)0.93 ML. EbW Fig. 1 U 4f spectra of U deposited on Pd at room temperature. U : Pd (AES) = (a)0.23, (b)1.05, (c) 1.92 and (d)5.96. 1 L (Langmuir)= Torr s-l. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 Table 1 Ratio of the CO 40 areas at high (10 L) and low (2 L) exposure, as a function of U coverage partial U coverage co 40 (10 L)/CO 40 (2 L) 0.00 2.00 0.54 1.oo 0.81 1.20 1.35 1.63 ~~ A decreased ratio indicates an increase in sticking probability of CO.that we are comparing values in the steeper initial part of the h = 2.94eV (a-*.) L111''1111111'IIII'IIIIIII'IIII II'I adsorption curve (Fig. 5). The experimental results show the opposite to be true. The ratio decreases, which indicates an 141210 8 6 4 2 0 E,IeV Fig. 3 UPS He I1 spectra after adsorption of 10 L CO on U/Pd at room temperature. U coverage: (a) 0.00, (b)0.37, (c) 0.56 and (6)0.93 ML. on which CO is not chemisorbed, and which thus blocks Pd adsorption sites. Second, the heat of chemisorption of CO on Pd may be changed by U so that at room temperature CO is more weakly adsorbed.Third, U may decrease the sticking probability of CO on Pd such that 10 L of CO would no longer be sufficient to saturate the surface. Let us discuss this in more detail. Fig. 4 shows the evolution of the CO 40 peak area with increasing U coverage. A steep initial decrease is followed by a more gradual one with increasing U dosage. The non-linearity observed can be explained either by block- ing of Pd adsorption sites consisting of several Pd atoms (ensemble effect) or by an additional change of the chemi- sorption properties of Pd by U. For a surface consisting of non-equivalent Pd adsorption sites, an overall decrease in chemisorption energy would leave fewer chemisorption sites capable of binding CO at room temperature. We will address this issue in more detail when discussing TD spectra.So far the experimental findings point to the first and second cases discussed above. To test the third case we compared the CO chemisorption signals at two different CO exposures (Table 1). A decrease in sticking probability would result in an increase in the ratio of the intensity of the CO signals (high- exposure signalflow-exposure signal) because the effective dosage (total dosage x sticking probability) decreased such 1.o h cn c.-5 0.8 G v 0.6 2 2 0.4 0 0 0.2 0'"""'"""""''"""""0 0.5 1.o 1.5 U coverage/ML Fig. 4 Evolution of the CO 4a/VB area ratio with U coverage after adsorption of 10 L CO at room temperature increasing sticking probability.Such effect has also been invoked as one factor for the promoting effect of alkali metalsI6 on the catalytic properties of metal surfaces. When the surface is covered by ca. 1 ML U the CO chemi-sorption signal is very weak and two new peaks appear at Eb = 4 and 6 eV (Fig. 3). We assign them to the 0 2p emis- sion from chemisorbed atomic oxygen and their presence shows that CO is adsorbed dissociatively at high U coverage. To enhance the weak oxygen features we subtracted the sub- strate background in the spectra of Fig. 6 until the emission at the E, disappeared. This procedure is somewhat delicate because, as we will discuss below, the valence band (VB) of U-Pd is modified by the presence of CO. For the two lower U coverages this procedure leaves a broad structure between 1 and 5 eV, part of which is due to changes in the shape of the VB after CO adsorption. The shoulder at 4 eV, which is also observed at high U coverage, might be due to chemi- sorbed oxygen because the 0 2p emisson of 0 on Pd appears at Eb = 4 eV." In this case some of the CO would dissociate even at low U coverage, but the resulting atomic oxygen would mainly interact with Pd and not with U because for 0 chemisorbed on U E, = 6 eV.6*'* This points to a strongly suppressed reactivity of U, which initially has a much higher affinity for 0 than Pd [AformH"(Pd0) = -20.4 kcal mol- '; A,,,,H"(UO,) = -257 kcal mol-1].'9 At a nominal U coverage of 1.3 ML (Fig.6), the 6 eV emission actually appears now as 0 bonded to U.It would have been inter- esting to study the evolution of the 0 1s emission as a func- tion of U coverage, because this would have allowed us to distinguish dissociated from molecular CO unambiguously. CO effective dosage Fig. 5 Comparison of two intensity ratios (high CO dosageflow CO dosage) for low (lz/li) and high (h,/h,) sticking probabilities. A decreased sticking probability results in an increase of the ratio (h,/h, -= 4/11). Fig. 6 UPS He I1 spectra after adsorption of 10 L CO on U/Pd at room temperature. The metallic background has been subtracted. U coverage: (a)0.5, (b) 0.8 and (c) 1.3 ML. However, the Pd 3p3,2 core-level line superimposes on the 0 1s line and, because it undergoes core-level shifts after U deposition and reaction, cannot be easily removed by a sub- traction or a curve-fitting procedure.CO adsorption on pure Pd leads to a strong suppression of the emission at the EF [Fig. 2(a) and 3(a)].This has been attributed to the fact that the electrons in the upper Pd 4d band (the antibonding part) participate in the CO-Pd bonding.20 Consequently a high DOS at the EF favours strong CO-Pd bonding. With increasing U concentration the VB becomes less affected by CO adsorption as seen by comparing the corresponding spectra in Fig. 2 and 3(a)-(d). This is partially due to a decreased CO concentration with higher U coverage. However, a more quantitative analysis of the VB shows that the decrease in concentration alone is not sufficient to explain this evolution.We measured the varia- tion of the intensity at the Fermi level normalized to the VB area between 0 and 6 eV (Table 2). A loss in intensity after CO adsorption is only observed for pure Pd. At low U cover- age there is no change and at high U coverage the intensity actually increases. This latter effect may be understood in terms of breaking U-Pd bonds by CO. As we discussed above, U-Pd bonding itself results in a decrease of the emis- sion at EF. Direct U-(C, 0)bonding, which exists at high U coverage, results in a breaking of the U-Pd bonds and the local DOS at EF on the Pd atoms increases again; at high concentration U indeed reacts with CO. We conclude that the decreased sensitivity of the Pd 4d band to CO adsorption after U coverage points to a reduced CO-Pd interaction, which would be in good agreement with a decreased amount of CO stable on the surface at room temperature.Table 2 Variation of the DOS at E, with CO adsorption, as a func- tion of U coverage normalized co partial U coverage coverage (CO/CO,,) i, i, (il -iz)/(il + i,) 0.00 1.00 0.032 0.0206 0.76 0.54 0.35 0.0107 0.0107 0.00 0.81 0.25 0.00874 0.00847 0.06 1.35 0.1 1 0.00544 0.00606 -0.22 The intensity at E, normalized to the VB area, is measured before and after 10 L CO adsorption (il and i,, respectively). The normal- ized difference (il -iz)/(il + i,)decreases with increasing U coverage. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 UPS Study of CO Adsorption at Low Temperature In the previous section we saw that U deposition on Pd inhi- bits the CO adsorption, either by simply blocking the Pd adsorption sites or by changing the chemisorption properties of Pd.We performed low-temperature adsorption of CO on U/Pd surfaces, followed by gradual heating of the surface to obtain further information on this phenomenon. We studied CO adsorption for U coverages below and above 1 ML. U coverage was calculated assuming that all U atoms stay on the surface even for full coverage which, as discussed above, is not entirely true above 1 ML because some of the U diffuses into the near-surface region. Thus the U coverage of 1.5 describes a near-surface interdiffusion layer which still con- tains some Pd atoms. The UPS study shows that even for such high coverages CO is chemisorbed on the surface at low temperatures (Fig.7 and 8), but with increasing U coverage the CO signal becomes attenuated as shown in Fig. 7, where all three curves have been normalized to the U-Pd VB. This normalization procedure actually tends to suppress the CO emissions because, with increasing U concentration, the cross-section of the VB increases because it becomes more like the U VB cross-section [5f3 (6d7~)~], which is larger than that of the Pd 4d.21 Thus the amount of CO chemisorbed on U/Pd is slightly higher than suggested by Fig. 7. It is remarkable that CO is chemisorbed associatively on a surface that is nominally completely covered by U. It may be that at these temperatures low-reactivity U atoms may act as adsorption sites, but we have also to consider that this effect may be produced by some Pd adsorption sites left in the surface layer or by an overestimation of the U coverage (because of uncertainties as to the information depth).Notice that the 40 for 1.24 ML U is enhanced when compared to the ln/5a emission. This is due to some physisorbed CO whose 1n/5a emission superimposes on the 4a line of CO chemi-sorbed on U. Physisorbed CO was not observed at lower U dosages and under the experimental conditions used. The heat of chemisorption of CO decreases with increasing U concentration. This is shown by the disappearance of the CO signal when the surface is warmed up: for a nominal U coverage of 0.95 ML heating to 0°C still leaves a CO signal [Fig.8(a)], while for a U coverage of 1.5 ML a complete disappearance results at this temperature [Fig. 8(b)]. Further- more, the difference spectra show that for a U coverage of 0.95 ML CO is desorbed from the surface with annealing, while for a coverage of 1.5 ML the increase of the 0 2p signal with annealing shows that at least some of the CO disso-'Is 1lI~1~~~1~~~1~~~1~~~l~~,l,,,l 16 12 a 4 0 E,IeV Fig. 7 UPS He I1 spectra after adsorption of 10 L CO on U/Pd at -165"C. U coverage: (a)0.00, (b) 0.66 and (c) 1.24 ML. J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 -8 "1,',1'''1'''"''1''' ,,, ,,T A I"'I"'I"'I"'I"'I"'I"' B Fig, 8 UPS He I1 spectra for adsorption of 10 L CO on U/Pd at -165 "C and subsequent annealing to room temperature.A, U coverage <1.0;(a) -150, (b) -100, (c) -50, (d)0 and (e)50 "C. B, U coverage > 1.0; (a) -165, (b) -150, (c) -100, (d) -50 and (e) 0 "C. The metallic background has been subtracted. ciates. This indicates that the reactivity of U increases with coverage. TD Study and CO Adsorption at High Temperature From the preceding discussion it is clear that CO chemisorp-tion on Pd is affected by the presence of U, which does more than simply block Pd adsorption sites. To get information on the heat of adsorption of CO on U/Pd we performed TDS experiments. We followed the evolution of the chemisorbed CO (CO 40 + 5a/ln) signal area when heating the sample on which we previously had adsorbed 2 L CO at -165 "C (Fig.9). Heating results in the sudden decrease of the CO signal. As we discussed above, this could be attributed (z priori either to CO desorption, at low U coverage, or to CO dissociation, at high U coverage. CO dissociation becomes the most likely process with increasing U concentration, because CO does not dissociate on pure Pd while on pure U it spontaneously decomposes. However, the U coverage used in the desorption experiment was below the threshold for CO dissociation, which lies between 0.95 and 1.5 ML nominal coverage [Fig. 8(a) and (b)].Thus we expect CO desorption rather than CO dissociation. For pure Pd we found a CO desorption temperature of 165°C(Fig. 9), which is in agreement with the literature.15 ~~~~~~~~~~~~~~~~~~'~~~~~~~~~~ -100 0 100 200 300 T/T Fig.9 TD spectra after adsorption of 10 L CO on U/Pd at -165°C. U coverage: (a) 0.00, (b) 0.25, (c) 0.75 and (d) 1.20 ML. Initial coverage, 2 L. Below this temperature the CO signal stays constant with increasing temperature, which shows CO to be adsorbed with one well defined energy, at least for a dosage of 2 L. With increasing U concentration the desorption temperature decreases gradually, which may be explained either by a geo- metrical effect, i.e. blocking of stable adsorption sites by U, or by a change of the overall chemisorption properties of Pd by U, which would be an electronic effect. The geometrical model assumes that there are different chemisorption sites at the surface, higher-coordinated ones with high CO chemi-sorption energy and lower-coordinated ones, onto which CO bonds less strongly.Partial coverage of the surface with U would result in a faster decrease of higher-coordinated adsorption sites than of the lower ones, which remain acces- sible to CO. CO would therefore be chemisorbed more weakly and hence the desorption energy is decreased. This model has been used to explain the decreased desorption temperature of CO on Au/Pd and Cu/Pd.22 However, it was shown that even for blocking 90% of the surface by Au the desorption temperature decreases only by 40°C. In our case covering the surface with U results in a decrease of the desorption temperature by up to 180"C, while the strong initial CO signal shows that more than 10% of the surface is available for CO bonding.Such a large decrease in the desorption temperature cannot be explained by a purely geo- metrical model and, therefore, indicates a change in the adsorption properties, probably the heat of chemisorption, of Pd by U. U indeed changes the electronic structure of Pd and, as we will argue below, this could very well modify its chemisorption properties. While the major effect of U is to decrease the desorption energy of CO on Pd, we also observed a secondary, opposite effect. When exposing a Pd surface, which was covered by 1.0 ML of U, to CO at 100°C part of the CO dissociates, as shown by the increase of the symmetrical 0 2p signal at 6 eV, indicating the formation of UO.Dissociation takes place because the U surface concentration is above the critical threshold coverage of ca. 0.9, above which U reactivity is high enough to induce CO dissociation; in fact, even if CO dissociation is an activated process it still occurs when CO is adsorbed at elevated temperatures. However, part of the CO does not dissociate but is adsorbed associatively. It is strongly bonded and heating even to 300°C does not result in its desorption, while on pure Pd desorption takes place below 200°C. This adsorbed CO does not change the desorption temperature for the majority of the surface CO in low- temperature adsorption-heating cycles, as discussed above Fig. 10 UPS He I1 study of high-temperature adsorption of CO.The nominal U surface coverage was 1.0. Tds= (a)and (c)-(e), -165 and (b) 100°C. Annealing temperature = (a)-(c), 100;(d)200 and (e) 300 "C. (Fig. 9), but the CO signal no longer drops to zero (Fig. 10). The C Is Ebin XPS is 287 eV, which is typical for chemi- sorbed CO. The signal at 13 eV in UPS is unexplained at present. If it is not due to a surface impurity building up during the repeated adsorption experiments, it has to be a product of the reaction of CO and the surface. Discussion We will discuss the observed changes in the surface reactivity of Pd with U coverage. The two main issues we want to address are the evolution in reactivity of the U surface atoms and the changes in adsorption properties of Pd caused by U.Reactivity of the U Surface Atoms Bulk U metal reacts immediately with CO to form a surface oxycarbide,6 which is a solid solution of UO and UC, where the 6d7s electrons participate in U-0 and U-C bonding while the 5f electrons remain largely non-hybridized. When U is deposited on Pd it loses much of this reactivity. At a nominal U coverage of 0.95 ML annealing leads to CO desorption instead of CO dissociation [Fig. 8(a)], which shows that for this coverage the activation energy for CO dissociation is higher than the desorption energy. At higher U concentration the activation energy for dissociation decreases, becomes smaller than the desorption energy and CO dissociation predominates over CO desorption. The dis- sociation product seems to be a U-0 complex as shown by the Eb of the 0 2p which is the same as for 0 chemisorbed on U [6 eV in Fig.8(b)]and differs from 0 chemisorbed on Pd (4 eV).17 The decreased reactivity of U may be attributed to the interaction between U and Pd, which places the initially highly reactive 6d7s electrons in stable solid-state bonds. The strong bonding interaction between U and Pd actually results in the decrease of the DOS at EF for both of the U and Pd atoms. However, at higher U concentration the emission at EF comes back and is now due to U 5f states (Fig. 2). The electronic structure of the U surface atoms begins to resemble J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 that of U metal and thus the reactivity increases.This experi- ment does not allow us to decide whether the decreased reac- tivity at low U concentration is directly linked to the decreased DOS at the Fermi level of the U atoms or whether it is due solely to the dilution of U on Pd. However, the decreased reactivity of U in UNi, ,6 where the 5f electrons are still delocalized, points to the dilution and solid-state bonding of U as the main factor. In addition, the high DOS at EFis due to the 5f states, which participate only weakly in bonding. Therefore, we would expect that in this case the DOS at EF plays only a secondary role. The experiments performed yield conflicting evidence about the role of U atoms as adsorption sites for CO. At low temperature, even for almost full U coverage of the surface, chemisorption of CO was observed, even though with some- what decreased intensity.On the other hand, even at satura- tion coverage the TD spectra did not reveal the existence of several adsorption sites with differing adsorption energies, as one would expect when both Pd and U act as adsorption sites. Instead, the desorption energy seems to decrease contin- uously with increasing U concentration and there is only one desorption peak [shown as a step, because we measured it in integrated mode (Fig. 9)]. It could be that the U atoms reside immediately below the surface, allowing Pd sites to interact directly with CO, even at nominal full U coverage on the surface. Adsorptive Properties of Pd U strongly affects the chemisorption properties of the Pd surface.The magnitude of this effect cannot be explained by preferential blocking of chemisorption sites, as was shown by comparing the U/Pd to the Au/Pd and Cu/Pd systems. Instead, the heat of chemisorption is lowered by U. This may be an electronic effect because even at low concentration U adsorption results in a measurable decrease of the local DOS at E, on the Pd atoms. This quantity has been directly related to the heat of adsorption and presented in Table 2, where the intensity at EF changes less with CO adsorption when the surface is precovered by U. The driving force for the bonding between U and Pd is indeed the filling of the Pd 4d band, which becomes a closed shell showing less ten- dency for interaction with CO.On the other hand, it has been argued that the heat of chemisorption of CO on late transition metals, will decrease when alloying the metal with electropositive elements having no d electrons (Si), thus resulting in the decrease of the CO adsorption energy. Alloy- ing with an electropositive element containing d electrons (i.e. an early TM) leads, in addition, to the weakening of the C-0 bonding because of the interaction between the TM d states and the CO antibonding 2n*.23Thus U may have two effects: it may lower the heat of chemisorption through bonding with Pd and weaken the intramolecular CO bonding by direct interaction. Decrease in the heat of chemisorption is not a local effect but it occurs through the changed electronic structure of the surface. Therefore, the presence of U does not create several different adsorption sites (there is one exception to this, as discussed below), but the adsorptive properties of the surface as a whole are changed.However, after high-temperature adsorption of CO we observe a secondary effect, which results in an increase in adsorption energy for some of the CO. Small amounts of CO stay on Pd at 300°C while on pure Pd it is desorbed below 200°C. Doping Pd with conventional promoters (K)also leads to strongly bonded C024 which is stable at high tem- peratures. Therefore, we think that by the high-temperature reaction between CO and the surface, some of the surface U is transformed into a promoter. Above a critical surface con- J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 1291 centration U reacts with CO to UO (and UC), i.e. it is par- tially oxidized to U2+which could have similar properties to alkaline-earth-metal cations (K'). This would also explain why after high-temperature reaction we observe two chemi- cally different chemisorption sites, a majority of sites whose chemisorption energy decreases with increasing U concentra-tion and a minority population of adsorption sites with strongly increased chemisorption energy. 3 4 5 6 7 T. E. Madey, J. T. Yates, Jr., D. R. Sandstrom and R. J. H. Voorhoeve, in Treatise on Solid State Chemistry, ed. N. B. Hannay, Plenum Press, New York, 1976, vol. 6B, p. 1. Z. Knor, in Catalysis (Specialist Periodical Report), The Royal Society of Chemistry, London, 1985, vol. 7, p.148. T. Gouder, Ph.D.Thesis, Namur, Belgium, 1987. T. Gouder, C. A. Colmenares, J. R. Naegele, J. C. Spirlet and J. Verbist, Surf:Sci., 1992,264, 354. R. Hauert, P. Oelhafen, R. Schlogel and H-J. Guntherodt, Surf Sci., 1985, 160, L493. 8 J. N. Andersen, J. Onsgaard, E. Zdansk, A. Nilsson and N. Mir- Summary We studied the CO adsorption on Pd surfaces covered by U. 9 10 tensson, Surf Sci., 1989,217, 127. T. H. Gouder and C. A. Colmenares, Surf Sci., 1993,295,241. M. P. Seah, J. Catal., 1979,57,450. U itself loses its reactivity at low concentrations where CO dissociation becomes an activated process. This passivation, which is probably due to the involvement of the highly reac- tive 6d7s electrons in stable solid-state bonds, provides the necessary condition for investigating the catalytic properties of actinides.On the other hand, U modifies the chemisorp- tion properties of the Pd surface in two different ways. First a change of the electronic structure of Pd, in particular the filling of the 4d band, results in a decreased chemisorption energy for most of the CO. This is a global effect that affects all surface Pd atoms. Second, after high-temperature reaction, some of the surface U becomes oxidized and starts behaving 11 12 13 14 15 16 17 18 19 M. P. Seah, J. Vac. Sci. Technol., 1980 17, 16. Y. Baer, H. R. Ott and K. Andres, Solid State Commun., 1980, 36,387. S. Doniach and M. Sunjic, J. Phys. C, 1970,3,284. Y. Baer, in Handbook on the Physics and Chemistry of the Actin- ides, ed. A. J. Freeman and G. H. Lander, Elsevier, Amsterdam, 1984, p. 271. B. Oral, Y.C. Lee and R. W. Vook, Appl. Surf Sci., 1990,44,65. H. P. Bonzel, Surf Sci. Rep., 1987,8,43. H. Conrad, G. Ertl, J. Kuppers and E. E. Latta, Surf. Sci., 1977, 65, 245. T. Gouder, C. A. Colmenares, J. R. Naegele and J. Verbist, Surf Sci., 1990,235,280. Handbook of Chemistry and Physics, CRC Press, Roca Baton, like a classical promoter increasing the chemisorption energy of co. 20 FL, 54th edn., 1973. R. Hauert, P. Oelhafen, R. Schlogel and H-J. Guntherodt, Solid State Commun., 1985,55, 583. 21 J. J. Yeh, and I Lindau, At. Data Nucl. Data Tables, 1985,32, 1. References 22 G. A. Kok, A. Noordermeer and B. Nieuwenhuys, Su$ Sci., 1 2 W. M. H. Sachtler and R. A. van Santen, Adv. Catal., 1977, 26, 69. R. A. van Santen, in Fundamental Aspects of Heterogeneous Catalysis Studied by Particle Beams, ed. H. H. Brongersma and 23 24 1985,153,505. R. Hauert, P. Oelhafen and H-J. Guntherodt, Sure Sci., 1989, 220, 341. A. Berko and F. Solymosi, Surf Sci., 1986, 171, L498. R. A. van Santen, NATO AS1 Series, Series B: Physics, Plenum Press, New York, 1991, vol. 265, p. 83. Paper 3/06 1 19A ;Received 13th October, 1993

ISSN:0956-5000    

DOI:10.1039/FT9949001285

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. CHEM. SOC. FARADAY TRANS., 1994, 90(9), 1293-1299 Spectroscopic Characterization of Magnesium Vanadate Catalysts Part 2.t-FTIR Study of the Surface Properties of Pure and Mixed-phase Powders Gianguido Ramis, Guido Busca and Vincenzo Lorenzelli lstituto di Chimica, Facolta di lngegneria, Universita P.le Kennedy, 1-16129 Genova, Italy The surface chemistry of the three stable monophasic powders magnesium orthovanadate Mg3(V0,), , magne-sium pyrovanadate Mg,V,O, and magnesium metavanadate MgV,O,, as well as of an amorphous oxide cata- lyst with an Mg : V atomic ratio of 3 : 2, has been investigated by FTIR spectroscopy. The bulk structure of the amorphous catalyst has been characterized by FT-FIR and Raman spectroscopies. The nature of the surface hydroxy groups, as well as the adsorption of molecular probes such as pyridine, acetonitrile and CO, , has been investigated.The samples show weak Lewis acidity and no Br~nsted acidity. Moreover, they show significant but not extreme basic character. They weakly adsorb oxygenate compounds like alcohols and ketones and interact only at high temperatures with alkanes and alkenes. These materials, active as alkane oxy- dehydrogenation catalysts, are much less reactive than other vanadia-based selective oxidation catalysts, like V,O,-TiO, and (VO),P,O, , towards organic molecules. This is attributed to the basic environment generated by the MgO component in catalysts belonging to the MgGV,O, system, which causes a parallel decrease of the Lewis acidity and of the oxidizing power of the active vanadium ions.Vanadium-based mixed oxides play a key role in industrial selective oxidation catalysis. Among others, the industrial catalysts for the synthesis of maleic anhydride from benzene oxidation are V,05-MOO, or v205-Wo3 mixed oxides,' those for the oxidation of butane to maleic anhydride are V205-P205mixed oxides,2 and those employed for the syn- thesis of phthalic anhydride from o-xylene oxidation belong to the V205-TiO, ~ystem.'.~ Recently, the VMgO system (V205-MgO) has been pro- posed as promising for the oxidative dehydrogenation of light alkanes to the corresponding light alkene~,~ a very attractive process according to the increased availability of light alkanes from natural gas. This catalytic system was pre-viously applied to styrene synthesis from ethylbenzene oxida- tive dehydr~genation.~ Since then, several research groups have investigated this reaction and this catalytic system.According to Kung and co-workers the active phase of the catalyst is the orthovanadate Mg3(V04)2 the high selec- ,476,7 tivity to dehydrogenation products being attributed to the absence of V-0-V bridges.8 Sheshan et al. correlated the IR V-0 stretching band of orthovanadates with the cata- lytic a~tivity.~ In contrast to Kung and co-workers, Volta and co-workers".' concluded that the actual alkane dehy- drogenation catalyst is the pyrovanadate Mg2V207. However, Bhattacharyya12 obtained good catalytic per-formances working with an MgO-V205 catalyst whose only X-ray diffraction (XRD)-detectable crystal phase was MgO.Burch and Crabb' concluded that the homogeneous reac- tion contributes to enhance the yield in alkene, and that no clear correlation exists between catalytic properties and crys- talline phase. On the other hand, the catalysts commonly used are p~lyphasic,~~'~~'~~'~ and an analysis of their phase composition including eventual XRD-undetectable species (like monolayers or amorphous phases) is not easy. The present paper summarizes some FTIR spectroscopic results on the surface and adsorption properties of different powders belonging to the MgO-V205 system. The aim is to try to correlate the surface properties of catalysts belonging to the MgO-V205 system with those of other vanadia-based catalysts, like V205-TiO, and V205-P205, which are active for the production of oxygenates instead of oxy-t Part 1: Ref.14. dehydrogenation compounds predominant on some Mg vanadates. Experimental The preparation and characterization of the pure MgO-V205 phase catalysts has been reported pre-viously.' O,' ',' An amorphous catalyst with an Mg: V atomic ratio of 3 :2 (8 m2 g- ') has been prepared by impreg- nation of fresh Mg(OH), with ammonium metavanadate fol- lowed by calcination at 673 K. Pure V205 (18 m2 g-') was from Degussa (Hanau, Germany). The IR spectra were recorded by a Nicolet Magna 750 Fourier-transform instrument. The skeletal spectra in the region above 400 cm-have been recorded with KBr pressed discs and with a KBr beamsplitter, while those in the far-IR region (400-50 cm-') have been recorded using the powder deposited on polyethylene discs, and with a 'solid substrate' beamsplitter.The IR spectra of the surface species have been recorded using pressed discs of the pure powders, in a heatable/liquid-nitrogen-cooledcell connected to a conven- tional gas-handling system. Removal of previously adsorbed species was carried out by outgassing, generally at 773 K. The FT-Raman spectrum has been collected with a Brucker RFS 100 instrument (Nd-YAG laser). Results and Discussion The experiments described here concern four MgVO cata- lysts. Three of them are the pure crystal phases Mg3(V0J2, Mg2V207 and MgV20,, whose catalytic activity has been reported previously.lo The fourth powder, here denoted the Mg vanadate catalyst, with an Mg :V atomic ratio of 3 : 2, appears amorphous to XRD analysis, although after calcina- tion at 873 K it crystallizes into a mixture of the three Mg vanadates. To obtain information on its structure, we investi- gated its FTIR, FT-FIR and FT-Raman spectra. Structural Characterization of the Amorphous Catalyst through Vibrational Spectroscopies The FTIR and FT-FIR spectra of the Mg vanadate catalyst are compared with those of V205 and MgO in Fig. 1. They can be discussed taking into account those of the pure phases 1294 Mg,(VO,), , Mg,V,O, and MgV,O, , discussed pre-vio~sly.'~The pair of sharp bands at 430 and 370 cm-', the shoulder near 550 cm-' and the broad band at 665 cm-' suggest that a phase very similar to MgV,O, is very likely to be present.On the other hand, the features at 980, 960 cm-', and at 337 cm-' show the presence of species similar to monoclinic Mg,V,07. The band near 845 cm-' can be due to the superimposition of the strong bands observed in the region 900-800 cm-' in all phases. On the other hand, note that the strong band apparent at 940 an-' and the shoulder at lo00 cm-' do not seem to belong to any of the three phases we characterized previously. Comparison with the spectrum of crystalline V205, also reported in Fig. 1, allows us to exclude the presence of this phase. On the other hand, amorphous vanadia also shows a strong V=O stretching band near 1020 cm- 1,15*16 like the crystalline compound, so its presence is also excluded.This seems relevant in view of the negative effect assigned to this phase with respect to cata- lytic activity in alkane oxidative dehydrogenation.' However, it seems very likely that MgO (in an amorphous form) is present too, although no definite evidence can be obtained by comparison of the spectra. Identification of the species present in the sample can also be attempted by analysing the 'first overtone' region, where IR-active combinations fall, typical for the crystalline phases.', In this region, the spectrum of the Mg vanadate catalyst [Fig. 2(b)] shows a sharp doublet at 1935, 1908 cm-', which provides definite evidence for the presence of Mg2V,0, .14 A strong, rather symmetric band at 1690 cm-' with a weak shoulder near 1790 cm-', in the region of the overtones of the terminal VO, stretchings, is also analogous to that observed for the divanadate Mg,V,O, .The triplet at 1220, 1150 and 1100 cm-' confirms the existence of V-0-V and/or V,O-V bridges. However, while the bands near 1220 and 1100 cm-' could be assigned to both 0.66 0.64 0.62 { 0.60 (0 s: 0.58; P (o 0.56 0.54 I 0.524 / ","'I"',"',' , I . ,I. ,. ;"'"''l'''l''' wavenumber/cm-' J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 J 2400.0 2200.0 2000.0 1800.0 1600.0 1400.0 1200.0 1000.0 waven u rnber/crn-Fig. 2 FTIR spectra of the Mg vanadate catalyst outgassed at room temperature (a) and at 773 K (b), and after following adsorption of pyridine (c) Mg,V,O, and MgV,O,, that at 1150 cm- ' does not belong to the pure phases we characterized in the overtone region.14 The definite absence of the band at 1370 cm-' and of the triplet nature of the band in the 1700 cm-' region seem to rule out the presence of microcrystalline Mg3(V0,), .Again, the combination modes of V205 are definitely absent. The Raman spectrum of the catalyst (Fig. 3) also clearly contains features assignable to the Mg,V,O, phase like the bands at 952, 900 and 633 cm-'. However, the strong band at 982 cm-' does not belong to any Mg vanadate species. Moreover, the presence of MgV,06, as well as of amorphous or crystalline V205 (whose Raman spectra are reported in ref.15), seems to be excluded. So, the IR spectra in the skeletal and first overtone region and the Raman spectrum suggest that this amorphous cata- lyst consists mainly of Mg,V,O, , although other struc-tures that should contain relatively short V-0 bonds (the IR bands at lo00 and 940 cm-' and the Raman band at 982 cm-') and V-0-V or V,O-V bridges (combination near 1150 cm-') should also be present. The presence of V205 particles is definitely excluded. Surface Characterization of MgVO Powders by IR Spectroscopy of Adsorbed Probe Molecules In Fig. 2 and 4, the IR spectra of the pure powder pressed discs of the Mg vanadate amorphous catalyst and of Mg,(VO,), , respectively, are reported before activation, and after activation and pyridine adsorption.In the region where the pure powder pressed discs partially transmit IR radiation ' '600' ' ' "400' " ' ' 0.2 ' wavenumber/cm-FTIR and FT-FIR spectra of (a) the Mg vanadate catalyst Fig. 1 (b)V,O, and(c) MgO Fig. 3 wavenumber/crn-' Raman spectrum of the Mg vanadate catalyst J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 + I 2400.0 2200.0 2000.0 1800.0 1600.0 1400.0 1200.0 1000.0 wavenumber/cm-' Fig. 4 FTIR spectra of Mg,(VO,), outgassed at room temperature (a)and at 773 K (b), and after following adsorption of pyridine (c). (d) Subtraction spectrum: (a)-(b). (e) Subtraction spectrum: (c) -(b). (above lo00 cm-I), together with the IR-active combinations of the skeletal modes, bands due to adsorbed species can also be found.Before activation [Fig. 2(a)], a very strong absorption is superimposed on the lattice combinations in the case of the Mg vanadate catalyst and this obscures the region between 1600 and 1400 cm-'. A similar but much weaker absorption is also found for Mg,(VO,), [Fig. qa)]. The subtraction spectrum [Fig. qd)] is dominated by bands at 1520, 1440 and 1080 cm-'. The same bands are also observed after adsorp- tion of CO, on the activated samples. Consequently, they can be assigned to carbonate species. The intensity of these bands compared with those of the bulk and of lattice combinations indicates that these species are far more abundant on the polyphasic amorphous catalyst than on Mg,(VO,), ,possibly this is related to the difference in surface area.These species need outgassing at near 573 K to be desorbed. Similar species, but in much smaller amounts, are also observable on the meta- and pyro-vanadate powders, and not at all on V205. Surface carbonates resistant to outgassing are typi- cally observed on metal oxide surfaces with basic character. They arise from the adsorption of atmospheric CO, and sometimes from the preparation method if carbonates or organic compounds (later burnt during calcination) are employed. However, if the oxide does not possess basic char- acter, these species, if present, are destroyed by room-temperature outgassing. These data suggest that the surface basicity of the catalysts increases with nominal MgO content, as expected. In spite of the poor transmittance of the samples in the region above 3000 cm-', due to radiation scattering, we also looked at the surface hydroxy groups of Mg vanadates.The spectra have a significant noise, but in all cases allow detec- tion of two weak absorption bands, one of which is very sharp near 3700 cm-' and the other slightly broader near 3650 cm-'. This is shown in Fig. 5 for the metavanadate and orthovanadate samples. According to our previous studies on Mg-containing spinels,I7 the band near 3700 cm-' is very typical for MgOH surface groups. On the other hand, a band near 3650 cm- has been found frequently on vanadia-based catalysts like V20,-Ti0, ,la and can also be envisaged on V205.19 This band can be assigned to VOH groups, in spite of the absence of Brarnsted acidity of these groups on Mg vanadates (see below).According to these assignments, which should both be taken as tentative, in all three pure-phase Mg wavenurnber/cm -Fig. 5 FTIR spectra (OH-stretching region) of MgV,O, (left) and Mg,(VO,), (right) outgassed at room temperature (a) and at 573 K (b) vanadates both Mg and V sites are actually located at the surface, although we cannot have information on a possible enrichment of one of them at the surface. The spectrum of the amorphous Mg vanadate catalyst in the OH stretching region (Fig. 6) strongly differs from those of the crystalline phases because of the presence of a strong, broad and split band with maxima at 3615 and 3510 cm-', which resists outgassing at 773 K.Stable species responsible for these features are very unusual on metal oxide powders with prevalently ionic character, but can be found in the spectra of predominantly covalent oxides like bulk V,O, l9 and WO,.,' In the present case, these species, not being associated with crystalline phases, should be related to amorphous phases containing covalent V-0 bonds. In Fig. 2(c) and 4(c) the spectra of the Mg vanadate cata- lyst and of Mg,(VO,), after pyridine adsorption are reported. In Fig. 7 the spectra of pyridine adsorbed on the Mg vana- date catalyst as well as on Mg,(VO,),, Mg2V207, MgV,O, and V205 are compared, after subtraction of the spectra of the corresponding adsorbents. The spectrum of pyridine adsorbed on pure V205 shows the presence of bands assign-ed to pyridine molecules coordinated on Lewis acid sites, 4.0 3.9 3.8 3.7 Q) 3.6 me g 3.5 a 3.4 3.3 3.2 3.1 -4 3800 3600 3400 3200 3000 wavenumber/cm-Fig.6 FTIR spectra (OH-stretching region) of the Mg vanadate catalyst outgassed at 673 K (a)and 773 K(b) 0) Cm fV2 1700 ' 1500 1700 1500 1700 1500 1300 wavenumber/crn-' Fig. 7 FTIR spectra of pyridine adsorbed on V,O, (a), Mg,V,O, (b),Mg,V,O, (c) and Mg,(V04), (d) and the Mg vanadate catalyst (e).For each compound, the upper and lower spectrum are relative to outgassing at room temperature and at 373 K, respectively, after pyridine adsorption. The last spectrum is reported only for V,O,.characterized by the sharp 8a mode shifted to 1608 cm-' (with respect to the liquid-phase value of 1583 cm-') and the 19b mode shifted to 1447 cm-' (from 1439 cm-' in the liquid). However, the bands at 1640, 1635 and 1535 cm-' also show the presence of pyridinium ions, evidence of the Brernsted acidity of this surface." The copresence of Brernsted and Lewis sites is also typical of other vanadia-containing catalysts like V,O,-TiO, catalysts18 and V~o~-P20~ cata-lysts?' On all MgVO catalysts only molecularly chemisorbed pyridine species are detected, bands due to pyridinium ions not being observed at all. No bands are observed near 1630 and 1530 cm-', so no significant Brernsted acidity is present on the Mg vanadate surfaces.Moreover, the positions of the bands of chemisorbed pyridine are slightly but definitely shifted to lower wavenumbers, with respect to the position observed on vanadia. We observe the 8a mode at 1605 cm- ' and the 19b mode at 1444 cm-l in all four Mg vanadate powders, without any marked difference among them. This means that the Lewis acid sites detected on Mg vanadates are weaker than those of vanadia. Comparison with the data obtained on V,O,-TiO, catalysts" and on v,o,-P,o, catalysts2' shows that Mg vanadates are the weakest solid acids of this series, as far as both the Lewis and Brernsted acid sites are concerned. This can be associated to the structural modification of the vanadium centres in these catalysts as a result of the presence of the basic component MgO, as dis- cussed previou~ly.'~ Identification of the Lewis sites as V ions arises from com- parison of the spectra of pyridine adsorbed on V,O,, Mg vanadates and MgO.In the last case, pyridine is adsorbed very weakly, with the 8a mode weakly shifted to higher wave- numbers (1595 cm-'). On the other hand, note that the struc- ture of Mg,(VO,), with its nearly cubic close-packed array of oxygen atoms with Mg in octahedral coordination and vana- dium in tetrahedral coordination is, according to Krishna- makhari and Calve,,, closely related to that of a cation deficient inverse spinel. It can be viewed as consisting of deficient spinel slabs related by antiphase boundarie~.~, We can consequently compare the surface properties of Mg3(V04), with those of the inverse spinel MgFe20, and those of the normal spinels MgAl,O, and MgCr,O,, investi-gated previou~ly.'~ From comparison of the data arising from pyridine adsorbed on these compounds, as well as on a number of other spinel-type compounds, we concluded that pyridine is adsorbed very weakly on Mg ions in unsaturated octahedral coordination, with the 8a mode below 1600 cm- '. We can consequently assign to pyridine on Mgz+ a band J.CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 near 1595 cm-' which we observe after contact with pyri- dine, but which disappears by simple outgassing at room temperature (Fig. 7), although an alternative assignment to pyridine molecules physisorbed on OHs cannot be definitely ruled out.The stronger sites responsible for pyridine coordi- nation resistant to outgassing, with the 8a band at 1605 cm-' are identified as vanadium cations exposed at the surface. As we noted previ~usly'~ the spectra of Mg3(VO4), pressed discs after activation show a very weak, sharp band at 1962 cm-l. This weak band is evident in Fig. 3(b), and disappears when basic molecules, like pyridine, are adsorbed. Consequently, this sharp band is rather evident in the sub- traction spectra of adsorbed pyridine as a negative peak [Fig. qe)]. Although the transmittance of the catalyst is extremely low below lo00 cm-', we can envisage in the subtraction spectra of adsorbed pyridine a negative band also at 982 cm-'.We consequently propose that these bands are the first overtone (1962 cm-') and the fundamental (982 cm-') absorption band relative to a surface mode that is perturbed upon pyridine adsorption. This mode is very likely to be the V-0 stretching of a surface vanadate ion acting as Lewis acid site for pyridine coordination. As discussed previo~sly,'~ the V-0 stretchings of bulk vanadate ions in Mg3(V04), fall in the region below 915 cm- '. However, these bands refer to V-0 bonds whose oxygen atom also coordinates two or three Mg ions, and which are vibrationally coupled. When a vanadate ion is exposed at the surface, one or more V-0 bonds can be free from cation coordination and their V-0 bond order can increase accordingly. Consequently, they can produce a surface V-0 stretching mode with a higher fre- quency than the bulk modes. These modes have already been found by us for different V,O,-containing catalytic systems, like V,O,-TiO, , where they fall at 1035 cm-' (fundamental) and 2045 cm-' (first ando~ertone),'~,~~on V,O, (1038 cm-', fundamental"). However, we showed that basic doping causes a strong shift of these modes on V20,-TiO, down to 1013, 1002 cm-' for the fundamental mode on K-V,O,-TiO, and down to 990 cm-' on Cs-V,O,-TiO, .24 Although coupling effects could also have a role, the values observed here (982 cm-', funda-mental; 1962 cm- ',overtone) can be taken as evidence for a weaker V-0 bond order for the surface vanadate species on Mg orthovanadate with respect to V205 and V,O,-TiO,. This is likely to be associated with the presence of tetrahedral vanadate species on the surface of Mg orthovanadate as well as in the bulk, in contrast to the square or trigonal-pyramidal coordination thought to be taken by vanadium ions on the surfaces of both V,O,-TiO, and V205.Similarly, the shift to lower wavenumbers of the V-0 surface stretching mode on V,O,-TiO, by basic doping, already reported,24 can be evi- dence of a progressive change from pyramidal to tetrahedral coordination for surface vanadium, due to the increased basicity of the oxide ligands. Moreover, the elasticity of the coordination of vanadium, from coordination four to coordi- nation six, also shown by the bulk structure of Mg vana- dates,', provides evidence for the ability of this cation to act as a medium-strong Lewis acid site.In Fig. 8 the spectra of acetonitrile adsorbed on Mg3(V0,), at room temperature are reported. Acetonitrile has been used previously by us as a basic probe for the acid sites of V205 and V20,-based catalyst^.'^*^' However, this molecule is also a probe for basic and nucleophilic sites; it tends to release hydrogen atoms from the methyl group giving the [CH,-CN]- anion upon attack of surface basic ions25-2 7 or to undergo hydrolysis due to attack of the nucleophilic sites on the electrophilic nitrile carbon at~rn.~'-~~On Mg orthovanadate, coordination of acetoni- trile on the Lewis acid sites certainly occurs, and is J. CHEM. SOC.FARADAY TRANS., 1994, VOL. 90 0.08441 1 0.0610 ! I;, -0.0557 2400.0 2000.0 1600.0 1200.0 wavenumber/cm-' Fig. 8 FTIR spectra of the surface species arising from the adsorp- tion of acetonitrile on Mg,(VO,), at room temperature (a) and after successive outgassing at room temperaure (b) responsible for the formation of the relatively sharp bands at 2268, 2295 cm-', due to CN triple-bond stretching and the 6(CH,) + v(CC) combination, which interact via a Fermi res- onance. The positions of these components confirm that the Lewis acid sites of Mg,(VO,), are weaker than those of V205, where these bands are detected at 2296, 2324 cm- ',as well as those of the V,05-TiO, and (VO),P,O, catalysts." These bands are shifted to higher wavenumbers from the liquid-phase values (2254, 2292 cm-') the stronger are the Lewis sites with which the molecule interacts. However, a strong band is also observed at 2211 cm-', whose intensity grows by increasing the contact time of the surface with acetonitrile.This band is similar to that observed on basic oxides like Mg0,26 Zn02' and CeO, 27 in the region 2200-2100 cm-', attributed to the CN stretching of coordinated [CH,-CNI- anions. This band is observed near 2050 cm-' for 'free' [CH,-CN]-ions in solution.28 This species, which is not observed on V,O, and other vanadia-based catalysts we investigated, is due to the exis- tence of basic sites on the surface of Mg,(VO,), which are able to abstract a proton from the methyl group of acetonitri- le.However, on a very basic surface like MgO, acetonitrile also undergoes hydrolysis of the CN bond, with the forma- tion of strong bands in the region 1800-1000 em-' due to amide, carboxylate and similar species. This has been shown for the acetonitrile-MgO interaction by Koubowetz et ~1.,'~ and has been confirmed by us. Analysis of this spectral region shows that this occurs, if at all, to a very small extent on Mg,(V0,)2. A comparatively weak band near 1700 cm- ' can be found [Fig. 8(a)], indicative of partial hydrolysis of the CN group. This indicates that only very few sufficiently strong nucleophilic sites are present, and this strongly dis- tinguishes the surface chemistry of Mg,(VO,), from that of MgO. IR Study of the Adsorption and Transformation of C,and C, Organic Compounds To obtain further information on the relationships between the acid-base and redox surface properties of Mg vanadates, as well as on the mechanism of their catalytic activity, their interaction with C, and C, organic compounds (i.e.the ketones acetone and methyl ethyl ketone, the alcohols isopro- pyl alcohol and sec-butyl alcohol, the alkenes propene and but-1-ene, butadiene and finally the alkanes propane and butane) has also been investigated. The results of these experiments will be compared with those obtained on vanadia-titania, and vanadyl pyrophosphate catalysts. '9, In Fig. 9 the spectra of acetone and isopropyl alcohol adsorbed on Mg,(VO,), are reported. Acetone is adsorbed weakly on Mg3(V0,), [Fig.9(b)] : the weakness of this inter- action is shown (i) by the poor vibrational perturbation of the molecule during this interaction; (ii) by its nearly complete desorption by outgassing at room temperature; (iii) by its sta- bility towards chemical transformation during adsorption. The CO stretching of adsorbed acetone is found at 1712 cm-' and the C-C-C stretching mode at 1231 cm- '. The position of these bands is intermediate between those of the gaseous molecule [1734 cm-', v(C10); 1215 cm-', V(CCC)~~]and those of the species adsorbed on V,O,-TiO, (1682, 1248 cm-I 29) showing that the interaction on the Mg vanadate is much weaker than that on V,O,-TiO,. The spectrum of isopropyl alcohol adsorbed on Mg,(VO,), is instead definitely similar to that observed on V,0,-Ti0,29 as well as on other metal oxides like pure TiO, .33 The sharp bands at 1468, 1385, 1368 and 1330 cm-' (the last one weak), are due to CH deformations of the iso- propyl moiety, while the very strong band with components at 1165, 1140 and 1125 cm-' is due to C-0 stretching coupled with C-C-C stretchings, typical of isopropoxy 1.75 1.70 1.65 1.60 m5 1.55B$ 1.50 1.45 1.40 1.35 11 0' '1700' '160C' 'l'5oo-'lbO0' '1300' '1200' '1'100' ' .' wavenumber/cm-' Fig. 9 FTIR spectra of the surface species arising from the adsorp- tion of isopropyl alcohol (a)and acetone (b) on Mg,(VO,), 1.701 n 1800 1700 1600 1500 1400 1300 1200 1100 wavenumber/cm-' Fig. 10 FTIR spectra of the surface species arising from the adsorp- tion of sec-butyl alcohol on the Mg vanadate catalyst, and successive evacuation at room temperature (a), 373 K (b), 453 K (c) and 523 K (4 groups.The weak band at 1300 cm-' is due to the deforma- tion of the OH groups of coordinatively adsorbed undis- sociated alcohol, present in small amounts, and desorbed by prolonged outgassing. The isopropoxy groups resist outgass- ing at room temperature and are progressively destroyed by heating under evacuation, disappearing near 473 K. During outgassing only traces of adsorbed acetone produced by its oxidative dehydrogenation are observed. In Fig. 10 the spectra of sec-butyl alcohol on the Mg vana- date catalyst are reported.Also in this case the adsorption of the alcohol is almost completely dissociative, as shown by the almost complete absence of the OH deformation mode near 1280 cm-', typical of coordinated sec-butyl alcohol, and by the presence of the strong C-O/C-C modes near 1100 crr-'. The bands near 1460 and 1380 cm-' are assigned to the asymmetric and symmetric deformations of methyl groups, that near 1420 cm-' to the scissoring mode of the methylene group and that at 1335 cm-' is the deformation mode of the methyne group. The alcoholates are progres- sively destroyed with the appearance of small amounts of methyl ethylketone near 450 K, evidenced by the sharp band near 1700 cm-'. These results show that polar molecules like the alcohols do react with the surface of Mg vanadates, giving rise to dis- sociative adsorption, while further dehydrogenation can also occur.However, the reactivity of this surface is much weaker than that of ~anadia-titania~'.~' where the alcohols are oxi- dized to ketones at 373 K, and the ketones are very strongly bonded, further transform to carboxylate species and finally completely decompose. This is mainly due to the stronger Lewis acidity of this surface than that of Mg vanadates, which causes stronger adsorption and greater perturbation upon adsorption. This conclusion, according also to our previous studies on the potassium doping of vanadia-based catalyst^,^, suggests that the Lewis acidity and oxidizing activity of V ions are strongly correlated, both being associated with electron with- drawal.Thus the oxidizing power of V ions is decreased when these centres are placed in a basic environment and this results in a lower oxidizing activity and, finally, in a lower catalytic activity in oxidation reactions. The weaker activity of Mg vanadates with respect to other vanadia-based catalysts is even more striking if the reactivity towards alkenes is considered. On the Mg vanadate catalyst, a surface interaction with but-l-ene is observed only near 523 K, with the formation of carboxylate species, characterized by bands at 1590 and 1420 cm-', due to the asymmetric and symmetric -CO, stretchings. However, these adsorbed car- boxylate species disappear by further heating at 673 K, owing to their complete combustion and/or the desorption of partly oxidized fragments.The same bands observed after contact with but-l-ene are also found after contact with buta-1,3- diene but at significantly lower temperatures (473 K). Closely similar behaviour is observed for the other powders belong- ing to the MgVO system. This behaviour strongly contrasts with that observed for vanadia-titania where alkenes are adsorbed strongly at room temperature producing alkoxy groups which are, in turn, easily dehydrogenated and oxi- di~ed.~'?~'The different behaviour of MgVO catalysts is associated with the lack of any surface Brransted acidity, in contrast to vanadia-titania where the Brransted sites easily protonate the alkene with the formation of alkoxy groups at room temperature or even lower.The reactivity of the Mg vanadates is also very different from that of vanadyl pyropho~phate~' where furan-like species and maleic anhy- dride are observed by interaction of C, alkenes near 523 K. The interaction with propane and butane does not give detectable adsorbed species on the Mg vanadates, although J. CHEM. SOC. FARADAY TRANS., 1994, VOL. 90 interaction occurs above ca. 623 K, according to the forma- tion of gas-phase carbon oxides and to the decrease of the transmittance of the sample probably due to reduction. This behaviour is justified by the fact that at the temperature at which this interaction occurs the resulting partially oxidized surface species are desorbed or are transformed further very rapidly. Conclusions The conclusions of the present paper can be summarized as follows: (i) A careful IR analysis in the skeletal and overtone region allows characterization of structurally polyphasic as well as microcrystalline MgO-V,O, catalysts.In particular, the presence of microcrystalline Mg2V20, and the absence of V20, on an XRD amorphous MgO-V,O, catalyst, have been established . (ii) MgO-V,O, samples are characterized by significant surface basicity, and tend to adsorb C02 in the form of car-bonate species. However, they are not as basic and nucleo- philic as MgO. (iii) All stable phases on MgO-V20, show weak surface Lewis acidity but a lack of any surface Brsnsted acidity, in contrast to pure vanadia and supported-vanadia catalysts. (iv) Weak Lewis acid sites are identified as coordinatively unsaturated surface V ions which are characterized on MgJVO,), by a surface V-0 bond much longer than the surface vanadyls on pure and supported vanadia.Therefore, they are probably essentially in a tetrahedral-like form, as in the bulk of these compounds, in contrast to the vanadylic form observed on supported-vanadia catalysts and vanadyl pyrophosphate. (v) On the surfaces of crystalline MgO-V,O, powders, OH groups bonded to both magnesium and vanadium are detect- able, showing that both cations are located at the surface, although surface enrichment cannot be excluded. For an amorphous powder, strongly bonded internal OHs are also evident.(vi) Mg vanadates show significant reactivity towards polar molecules such as alcohols, which are adsorbed in a disso- ciative way. (vii)Mg vanadates are very poorly reactive towards hydro- carbons like alkanes and alkenes. The weak interaction with alkenes, associated with the absence of Brsnsted acidity, can explain the ability of these powders to catalyse selectively the oxidative dehydrogenation of akanes, with limited alkene overoxidation. The role of the MgO component in these oxy-dehydrogenation catalysts seems to be essentially related to the lowering of the acidity (probably of both Brransted and Lewis type) and consequent- ly of the oxidizing ability of the vanadium oxide key com- ponent, which limits the successive transformation of the desired products (alkenes and butadiene) although it also strongly limits catalytic activity.This work has been supported by MURST (Rome). The authors thank D. 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ISSN:0956-5000    

DOI:10.1039/FT9949001293

出版商:RSC    

年代:1994

数据来源: RSC