Metal-Catalysed Reactions of Hydrocarbons / 12-Dehydrogenation of Alkanes

below 373 K, and above 473 K dehydrogenation without hydrogenolysis was observed: the distributions of deuterium atoms in reactant and product were similar, indicating that random loss of both kinds of atom occurred in the dehydrogenation steps. The transient response method has also been applied.47 Pt / KL zeolite was more active than Pt/SiO2, the addition of potassium to which lowered activity but improved thiotolerance.

Cycloalkane dehydrogenation has been regarded as the archetypal structure- insensitive reaction,48−52 but recent work has qualified this belief. Changes in the free surface area during reaction are due less to particle growth53 than to ‘carbon’ deposition; the character of the reaction was consequently found to change from sensitive to insensitive.10 This provides the clearest possible indication that the ‘carbon’ formation homogenises the surface, eliminating high activity sites, the concentration of which is size-dependent, and leaving only low activity sites that are common to all sizes. One other factor has been identified:53 high-temperature reduction of Pt/α-Al2O3 creates a strongly-held form of hydrogen, particularly on (or in?) very small particles, and this acts as a poison. Its removal by flushing with argon increased the TOF by about 40-fold, and this altered the form of structuresensitivity: after argon purging, highly dispersed catalysts were the more active. However, with freshly reduced catalyst of dispersion less than 50%, the effect of particle size on TOF might well be missed. There have been many other indications that strongly held hydrogen in platinum catalysts is harmful to their activity in hydrocarbon reactions.54

If, however, conditions are chosen appropriately, rates are a linear fraction of amounts of hydrogen chemisorbed: thus with platinum and rhodium on ceria, alumina and their mixtures, such a linear plot has been used55 as a means of estimating dispersion from activity measurement, this being a more sensitive and simple procedure than determining an adsorption isotherm. Thus after reduction at 1273 K, the dispersion of a platinum catalyst was estimated to be 3%.

Modification of ensemble size is achieved not only by introducing an inert or low activity element (see next section), but also by using a support or an oxidic component that can induce the Strong Metal-Support Interaction. Thus with Rh/SiO2 at 733 K, the specific rate for cyclohexane dehydrogenation was independent of reduction temperature, but reduction of Rh/Nd2O5 at 473 and 773 K gave rates that were respectively about 10 and 40 times smaller.56 The activation energy ( 105 kJ mol−1) was affected, but was smaller than for Rh/SiO2 (151 kJ mol−1). The use of niobia as support also suppressed simultaneous hydrogenolysis.22,57 Addition of manganese to Rh/SiO2 after calcinations gave the mixed oxide Rh2MnO4/SiO2, which on reduction at 573 K resulted in small rhodium particles modified by interaction with MnOx;58 however, after reduction at 473 K, high activity for converting cyclohexane was observed.

Although little work has been published on the other metals of Groups 8 to 10, some of them feature as one terminus of a bimetallic series to be considered

presently. In the reaction of cyclohexane in the presence of deuterium on nickel powder, deactivation above 373 K lowered activity 100-fold, without changing activation energy;33 but above 473 K, the products were benzene, n-hexane and (strangely) some toluene. Exchanged products declined as these other reactions set in, and under these conditions exchanged cyclohexane was not returned to the gas phase. Palladium on sepiolite or AlPO4-SiO2 caused disproportionation (process 10.B) at low temperature and dehydrogenation when the temperature was raised sufficiently to cause hydrogen to desorb.59 Pd/Al2O3 (D= 40%) was 20 times less active than Pt/Al2O3(D = 50%);60 this work was supported by calculations using SCF/CNDO methodology, which confirmed C––H bond breaking as the slow step. Methylcyclohexane dehydrogenation was four times faster on Pd(111) than on the (110) or (100) surfaces.61

12.3.3. Reaction on Bimetallic Catalysts4,62

There is extensive work to report on dehydrogenation of cycloalkanes on bimetallic catalysts, mainly in the supported form. Its motivation is quite clear: to minimise parasitic reactions such as ‘carbon’ deposition and hydrogenolysis, so as to have a catalytic system capable of working for long periods of time and at high selectivity.

Studies with single-crystal surfaces27,63 have helped to elucidate the modifying effect of tin on platinum catalysts, as indeed they have in the case of isobutane.27 Chemisorption of cyclohexene on Pt(111), Pt3Sn(111) and Pt2Sn(111) showed that its strength of adsorption decreased with increasing tin content, and that it was unable to form a di-σ bonded species on the last,63 although adjacent pairs of platinum atoms still existed. The LEED structures of these surfaces were shown in Figure 4.2. This tendency might help to explain the beneficial effect of tin in this reaction;64 ‘carbon’ deposition under reaction conditions was also suppressed. Deposition of tin atoms onto Pt(111) increased the TOF for cyclohexane dehydrogenation at 773 K, the maximum enhancement (×1.7) being at 20% tin coverage.27 Both these studies imply some electronic modification of platinum by tin.

Experiments with PtPb/Al2O365 and with PtSb/Al2O366 have shown that these additives have similar effects to those produced by tin.65 They decrease activity at low temperature by occupying active centre, but their beneficial effect is shown at high temperatures, where they diminish ‘carbon’ formation on the metal, and help to maintain activity. In trying to assess the role of the second component, attention must be given to the conditions under which the measurements are made.

The inclusion of tin in an Ir/Al2O3 catalyst (5wt.% of each) completely suppressed the hydrogenolysis of cyclohexane, which in its absence gave n-hexane as a major by-product;67 the rate at 526 K was however decreased, but the activation energy fell from 208 to 125 kJ mol−1. This remarkable and important result

seems to have escaped attention: such a catalyst might compete successfully with PtSn/Al2O3.

In the combination PtMo/SiO2, activity for cyclohexane dehydrogenation at 543 K decreased linearly with increasing molybdenum content;68 no significant difference in rate at 513 K was found between Pt/C and Pt3Zr/C catalysts.69 Rhenium is much more active than platinum for hydrogenolysis, but when a bimetallic catalyst containing them is pre-sulfided a ReSx species is formed that acts as a selective site-blocking agent, leaving small ensembles of free platinum atoms available for reaction. With Pt(111) the maximum rate of cyclohexane dehydrogenation occurred when θRe was 0.5.70 A thorough kinetic study of the dehydrogenation of methylcyclohexane on Pt/Al2O3 and on PtRe/Al2O3 presulfided showed that the latter was less active and had the higher activation energy (133 as against 196 kJ mol−1), the rate being unaffected by changing hydrogen pressure. The results were modelled using Hougen-Watson methodology, and thermodynamic parameters were derived.71 The activity of PtRe/Al2O3 at 843 K for the cyclohexane reaction only declined when the rhenium content exceeded about 50%.72,73 Sulfiding led to better selectivities.74

The platinum-ruthenium couple is well known to exhibit synergism in a number of catalytic processes, especially those of an electrochemical character, but the activity of ruthenium for hydrogenolysis greatly exceeds that of platinum.75 With PtRu/Al2O3, this was a linear function of composition in the cyclohexane-hydrogen reaction, but a maximum rate of dehydrogenation at 570 K was found when the surface contained about 55% platinum.76

Some very important conclusions have been obtained from work on the platinum-gold system.41,77 Deposition of gold atoms onto Pt(100), and of platinum atoms onto Au(100), both increased the rate of cyclohexane dehydrogenation at 373 K, the maximum rate being some five times greater than for Pt(100) alone.77 In the latter case, the maximum occurred after coating with enough gold atoms to form about a monolayer, but it appeared that island formation took place, leaving some small platinum ensembles uncovered. A similar enhancement was found with PtAu(111) surface alloys at 573 K;41 in both cases a small amount of cyclohexene was also detected.

It finally remains to note several studies in which metals of Groups 8 to 10 other than platinum have had their reactivity towards cyclohexane dehydrogenation moderated by inert additions. In a very famous paper,78 extensively cited, a series of nickel-copper powders were prepared and found to be homogeneous by XRD, magnetisation and hydrogen chemisorption measurements. The specific rate of cyclohexane dehydrogenation at 589 K was essentially constant the activation energy being about 220 kJ mol−1 for nickel contents between about 20 and 95%. This rate was a little greater than that for pure nickel, and these results were in stark contrast to those for ethane hydrogenolysis, the rate of which decreased dramatically as nickel content fell (see Chapter 13). They provide a telling demonstration that for

this reaction at least particle-size and ensemble-size sensitivities are equivalent. This is consistent with the view that a single active atom is sufficient to effect reaction. Nickel and palladium have also been beneficially promoted by addition of tin.79,80

This informative work on the nickel-copper system led to a number of related studies. A further important paper by Sinfelt showed that bulk mutual solubility was not an essential precondition for obtaining interaction between active and inactive metals in small particles. Thus silica-supported ruthenium and osmium were both modified by the presence of copper,34 but for the cyclohexane-hydrogen reaction it was only the rate of hydrogenolysis that was lowered. Benzene selectivities therefore rose to 95% at Ru:Cu or Os:Cu ratios of unity. The interaction of ruthenium and copper depends somewhat on the type of silica used, however.81 Benzene selectivity and rates were also increased by adding silver to rhodium, the RhAg/TiO2 catalyst at 573 K showing the extremely high selectivity of 99.6%.82

12.4. THE CHEMISORPTION OF HYDROGEN ON PLATINUM2

Clearly the strength of chemisorption of hydrogen is one of the factors determining activity in dehydrogenation, because if the atoms are reluctant to recombine and desorb the alkane will be unable to find the necessary free sites for its own adsorption: this effect is manifested in negative orders in hydrogen, and ultimately the reaction would grind to a halt. However, the requirement for vacant sites conflicts with the need to maintain a presence of hydrogen on the surface to counteract excessive dehydrogenation, and the consequential development of ‘carbon’ deposits. A catalytic system that is successful in practice, and stable in the long term will have to adopt a set of conditions in which these conflicting needs are reconciled.

Nevertheless in the steady state on Pt/Al2O3 the transfer of hydrogen atoms between the surface and hydrocarbon species formed from cyclohexane, and the desorption of benzene, have been shown to be fast compared to cyclohexane chemisorption and hydrogen desorption, which therefore control the overall rate. Now although our earlier discussion of hydrogen chemisorption on metals (Chapter 3) revealed nothing unusual about platinum, a careful consideration2 of the literature shows rates of desorption of hydrogen from platinum are uniquely slow, and the strength of adsorption unusually strong. This is shown by (i) very low values for the pre-exponential term associated with the desorption rate constant, and (ii) a very high value for the adsorption coefficient on Pt/SiO2 compared to those for PtRh/SiO2 bimetallic catalysts. No agreed explanation for these fundamentally important observations has been forthcoming, but it seems probable that an important contributing reason for the success of bimetallic catalysts containing platinum is that the additive causes the chemisorption of hydrogen to be weakened.

There are experimental results to support this assertion. First, where orders of reaction in hydrogen can be compared, they are less negative on bimetallic catalysts than on platinum alone.8,71 Secondly, the temperature-programmed desorption (TPD) profiles of platinum catalysts were changed by the presence of tin, which gave rise to a marked increase in the size of the low temperature (623 K) desorption peak. Corresponding experiments with single crystals however showed the opposite effect with Pt3Sn(111), and the same difference only with Pt2Sn(111). There is the additional possible complication with platinum catalysts of very strong held (dissolved?) hydrogen formed in high temperature treatment, the negative effect of which has already been noted: it is quite possible that this might form at temperatures commonly used for dehydrogenation (673–773K), but it seems not to be known whether tin or other additives prevent this.

The kinetic study of isobutane dehydrogenation mentioned above23 shows clearly that with Pt/Al2O3 at 900 K increasing the hydrogen/isobutane ratio favoured dehydrogenation over carbon formation, but the use of high ratios is uneconomic. Although at the same temperature the beneficial effect of hydrogen was much less marked with PtSn/Al2O3 because of the weaker adsorption of hydrogen; nevertheless it can operate at lower hydrogen:isobutane ratios since the form of ‘carbon’ deposit is changed, more going to the support and less to the metal. The possibility that the ‘carbon’ deposit may act as a reservoir for hydrogen in this and other reactions will be considered below.

Bewilderment is sometimes expressed that a catalysed reaction is able to bring a reacting system to that position of equilibrium that would hypothetically be reached without it, and that the catalyst is somehow able to anticipate the state of equilibrium that should pertain only to free molecules. There is indeed evidence in many systems that the products returned to the gas-phase are not in the proportions that meet thermodynamic expectations: this is seen most clearly in the hydrogenation of, for example, 2-butyne, where the less stable Z -2-butene is formed preferentially (see Chapter 9). With cyclohexene and hydrogen at about 473 K the catalyst has to decide whether to make benzene or cyclohexane, utilising chemisorbed species that differ greatly from their gas-phase counterparts. In these cases the initially formed products are fixed by factors relating to adsorbed intermediates, but repetition of the adsorption-reaction-desorption cycle eventually allows the proper position of equilibrium to be attained. In the case of the 2-butyne reaction, isomerisation of the 2-butenes has to await the near-total removal of the alkyne in order for them to gain access to the surface. When this can occur, the relative stabilities of the various conformations, determined by repulsive interactions between non-bonded atoms, allow reaction to proceed in the general sense of what equilibrium demands, although those interactions are not identical with those that operate in the free molecules. With cyclohexane dehydrogenation, the effect of rising temperature in decreasing the adsorption coefficients of hydrogen and of benzene enables reaction to go in the expected direction.

12.5.THE FORMATION, STRUCTURE, AND FUNCTION OF CARBONACEOUS DEPOSITS83,84

It may be helpful at this point to try to draw together some of the many threads concerning ‘carbon’ deposition that have appeared many times in the previous chapters: a comprehensive and unifying model is not yet available, and indeed it is doubtful if it ever will be, so many are the factors that determine the state of ‘carbon’ on metal surfaces. Nevertheless it is possible to make a few generalisations, and an attempt to do so is opportune now because the last section of this chapter concerns a constructive use of surface carbon to create useful products. The term ‘carbon’ will be used as an omnium gatherum for what has been variously named coke, acetylenic residue, carbonaceous deposit and probably other things as well. The following short survey may be amplified by reference to review articles.1,25,83,85

The form and quantity of ‘carbon’ existing on the surface of a metal catalyst depends inter alia upon the following variables: (i) the nature of the metal, (ii) its physical form, i.e. single crystal, powder or ‘black’, small supported particle etc., (iii) the nature of the support, if any, (iv) the type of hydrocarbon applied, (v) the presence of other molecules, especially hydrogen, and the hydrogen: hydrocarbon ratio, (vi) the time and especially the temperature of exposure. We may briefly consider the importance of each of these factors.

Within Groups 8 to 10, the base metals (Fe, Co, Ni) are distinctly more prone to form ‘carbon’ than the noble metals: metals to the left of these Groups (except perhaps manganese) are also very susceptible, so there is qualitative correlation between a metal’s propensity to form stoichiometric or non-stoichiometric carbides, the strength of C––M bonds that it can form at the surface, and its tendency to deposit ‘carbon’ under the appropriate conditions. As we have seen, lowering the size of the ensemble of active atoms including an inert partner (e.g. PtSn, PtAu, RhAg) does not always reduce ‘carbon’ formation, as the additive may have other undesirable effects, but in general this move is usually helpful. This may be because (in the case of tin) it weakens the bonds of hydrocarbon species to the surface, so assisting their migration away from the metal. The same effect is produced by sulfided platinum-rhenium.70 It may be (in the case of tin) that it weakens the bonds of hydrocarbon species to the surface, so assisting their migration away from the metal.86

Somorjai and his colleagues have developed a model87 for the states of ‘carbon’ on a platinum surface containing steps and kinks, in which much of the surface was obscured by a ‘carbonaceous overlayer’ with islands of 3D ‘carbon’, leaving only a few single atoms or pairs at steps uncovered. It was felt that the higher activity of sites at steps would cause hydrogen if present to break C––M bonds. If this is so, then very small metal particles that expose only atoms of low coordination number should be more resistant to ‘carbon’ deposition than larger particles, powders or macroscopic forms. Quantitative evidence on a particle-size effect is

difficult to find: while supported metals are not immune to ‘carbon’ formation, highly reproducible activity can be found, even in such an unpromising system as ethyne hydrogenation on nickel/pumice.88 The protracted use of macroscopic forms over a period of days is rarely if ever attempted, so exact comparison with supported metals is impossible, but intuitively one feels they would not last the course.

The process of ‘carbon’ deposition begins with the removal from the chemisorbed hydrocarbon of more hydrogen atoms than are strictly necessary to achieve the intended process. As we have seen in Chapter 4 (Figure 4.2), the progressive dehydrogenation of ethene proceeds through the sequence83,84

ethylidyne = vinyl → vinylidene → ethyne → ethynyl

We shall see in the next chapter, however, that in alkane hydrogenolysis it is probably necessary to remove two or even three hydrogen atoms to form species in which the C––C bond is so strained that it breaks, before the process of hydrogen addition can start. Thus species that are irrelevant to and toxic towards reactions such as alkane exchange and alkene hydrogenation may be the necessary intermediates for reactions such as hydrogenolysis, where they become reactive at higher temperature. The Pt––C bond is essentially non-polar, since NEXAFS studies of deactivated Pt/Al2O3 indicate the absence of electron transfer.

‘Carbon’ formation can then proceed in various ways. The di-carbon species shown above may break down to mono-carbon species then polymerise, or polymerisation may occur without this. Processes in which free radicals occur, or acidity in the catalyst, both encourage this, and (as noted above) catalysts for dehydrogenation must be neutral or basic to prevent acid-catalysed reactions of the alkene. The more dehydrogenated the reactant hydrocarbon, the greater the tendency to ‘carbon’ formation: alkadienes89 and aromatics90 are notorious in this respect. With alkanes, increase in molar mass (i.e. the number of C––H bonds) also assists degenerate events.

Because of its importance and frequent occurrence, the process of ‘carbon’ formation and its structure have been widely studied, and many physical techniques83 (including recently positron-emission tomography91) have been deployed. Of these, temperature-programmed methods (oxidation, TPO; reaction with hydrogen, TPRe) are the simplest and most informative. TPO can distinguish between ‘carbon’ on the metal, which is relatively easily oxidised, from ‘carbon’ on the support, which is less reactive. Admixture of the sample with a Pd/SiO2 catalyst ensures that the effluent contains only carbon dioxide, and no monoxide.92 The use of the TPRe strictly requires estimation of the methane (and possibly other alkanes) that emerges, since because the H/C ratio in the ‘carbon’ is unknown, the

amount of hydrogen consumed is not an accurate indicator of the carbon content. The ratio varies widely and depends upon catalyst composition2 and especially temperature, since the end products are either amorphous carbon, or (frequently) graphite, or even above 870 K a diamond-like sp3 form that is difficult to oxidise or reduce.93A NEXAFS study has revealed no electron transfer between platinum and carbon, so there can be no electronic modification of neighbouring free sites.94

Many different manifestations of ‘carbon’ have been recognised, and an attempt is made to show their interrelations in Scheme 10.2. ‘Carbon’ may either dissolve into the metal, forming some kind of carbide, or it may cover some or all of the metal surface, or migrate to the support, where it can deactivate by blocking the pores of a microporous catalyst; or even encapsulating the whole particle; or most interestingly it may grow away from the surface as filaments. These were formerly called the vermicular form, but are now recognised to be tubular forms (nanotubes) having structures related to buckminsterfullerene;95 they grow particularly well on base metals, and usually carry a small metal particle at their head. They are now prepared on a significant scale by the reaction of ethyne with a base metal at high temperature, and are the subject of intensive study because of the uses they may prove to have in solid-state devices.

From the academic point of view, ‘carbon’ formation is regarded as an unmitigated nuisance by those trying to study hydrocarbon reactions and to obtain quantitative information on them for purposes of mathematical modelling. In practice one either has to work with a catalyst in its stable but highly deactivated state,96 or to use very short reaction periods interspersed with cleansing times when only hydrogen is passed.97 Even so, frequent recourse must be had to ‘standard’ conditions, to ensure constancy of activity, or to monitor deactivation should it occur: in this case, bracketing with ‘standard’ periods can still lead to usable results. The pulse reaction mode, in which hydrocarbon pulses are injected into a hydrogen stream, is unsuitable for precise work.

Perhaps the most contentions and often-considered aspect of ‘carbon’ is its possible actual participation in catalysed processes. The idea originated with the suggestion by Thomson and Webb98,99 that alkene hydrogenation proceeded on top of a ‘carbon’ layer, utilising hydrogen atoms associated with it, and accounting for the structure-insensitivity and constant activation energy shown by this process. It is certainly true (or as true as anything can be in catalysis) that exposure of metal surfaces to ethene leads to the rapid formation of a layer of ethylidyne species (8),100 but the idea originally canvassed that hydrogenation occurred through the ethylidyne ethylidene equilibration could not be sustained84 (see Chapter 7). Ethylidyne does not however totally prevent hydrogen chemisorption, and indeed it now appears that it actually promotes di-σ ethene chemisorption on Pd(111).101 ‘Carbon’ inhibits the hydrogenation of alicyclic molecules, but has lesser effect on the reactions of alkenes. Unfortunately some obvious and critical experiments with isotopic labels have never been performed, nor has the concept of ‘carbon’ acting

as a hydrogen reservoir ever been quantitatively modelled: only in the case of benzene hydrogenation has it been formally incorporated in a reaction mechanism (Chapter 10). If indeed it were the vehicle by which hydrogenation occurred, it would be hard to explain the observed orders of reaction, and particularly the characteristic patterns of behaviour that distinguish the metals of Groups 8 to 10.

Little attempt seems to have been made to estimate the number of ‘free’ surface metal atoms in coked catalysts, and hence to find TOFs, assuming these to be the seat of the residual activity. While the use of hydrogen chemisorption might be considered risky, that of carbon monoxide ought to be suitable. Based on the loss of its IR intensity, the active metal area of Pt/Al2O3 used for n-heptane reforming was only 8% of its initial value, but its extent of adsorption slowly increased as it displaced some of the ‘carbon’.102

The spectator of research on ‘carbon’ might well sympathise with Dogberry’s belief that They that touch pitch will be defiled.

12.6. THE HOMOLOGATION OF METHANE85

Every cloud has a silvery lining, and there is one possibly useful way of employing the process whereby methane after dissociation at a metal surface enters a C––C bond-forming reaction. Although methane is abundantly available, it is of limited value except as fuel or as a feedstock for steam-reforming; the concept of its homologation into high alkanes of great utility was therefore a most attractive one, and its activation under reducing conditions promised better success than has ever attended its oxidative dehydrogenation to methanal or methanol. The reaction was therefore intensively studied through the 1990s and a considerable literature resulted.

Although the first work, and much subsequent effort, employed platinum catalysts103−109 (especially EUROPT-1, 6.3% Pt/SiO2), it appeared that better results would be obtained with metals that had a reputation for extensive C––C bond formation, starting from C1 species as in Fischer-Tropsch synthesis:110 attention was therefore focused on cobalt105,107,110−114 and ruthenium.71,105,106,110,115 Conventional supports have been used, but zeolites have also been investigated: the problem of fully reducing cobalt ions on alumina or in zeolite has been overcome by including a more easily reducible metal (e.g. ruthenium110,113) which assists the reduction, probably through hydrogen spillover. Other metals (e.g. Rh,116 Pd,117 Cu118) have been looked at in a cursory manner.

The amount of methane chemisorbed may be estimated by the amount of hydrogen released, or better by TPO of the retained ‘carbon’. Chemisorption at high temperature leads to several distinguishable forms of ‘carbon’ that differ in reactivity,116 but the composition of mono-carbon species CHx formed at lower temperatures (573–723 K) can be assessed by the composition of deuteromethanes

formed by their reaction with deuterium.109 The best chance of C––C bond formation is when x is two or three, but in general the procedure has been to decompose methane at high temperature (up to 973 K) and to hydrogenate the ‘carbon’ at lower temperature to avoid the hydrogenolysis of any ethane or propane that may be formed. No conditions seem to have been found under which useful yields of higher alkanes have been obtained, and aspirations to found a major chemical process have been disappointed. Interest in this subject has therefore waned.119 The few references cited should lead any interested reader towards other relevant papers.

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