Hydrogen Storage in Liquid Organic Hydride: Producing Hydrogen

REVIEW pubs.acs.org/EF

Hydrogen Storage in Liquid Organic Hydride: Producing Hydrogen Catalytically from Methylcyclohexane Faisal Alhumaidan,*,† David Cresswell,‡ and Arthur Garforth‡ † ‡

Petroleum Research and Studies Center, Kuwait Institute for Scientific Research, Post Office Box 24885, Safat 13109, Kuwait Environmental Technology Center, School of Chemical Engineering and Analytical Science, The University of Manchester, Post Office Box 88, Sackville Street, Manchester M60 1QD, United Kingdom ABSTRACT: Hydrogen storage for stationary and mobile applications is an expanding research topic. One of the more promising hydrogen storage techniques relies on the reversibility and high selectivity of liquid organic hydrides, in particular, methylcyclohexane (MCH). The use of liquid organic hydrides in hydrogen storage also provides high gravimetric and volumetric hydrogen density, low potential risk, and low capital investment because it is largely compatible with the current transport infrastructure. Despite its technical, economical, and environmental advantages, the concept of hydrogen storage in liquid organic carriers has not been commercially established because of technical limitations related to the amount of energy required to extract the hydrogen from liquid organic hydride and the insufficient stability of the dehydrogenation catalyst. This paper provides a review for the effort that has been directed toward the development of this concept over the past few decades and mainly focuses on the catalytic production of hydrogen from MCH. The topics that have been covered are the kinetics of MCH dehydrogenation over Pt/Al2O3 and Pt
Re/ Al2O3 catalysts, the kinetics of catalyst deactivation, the thermodynamic equilibrium in MCH dehydrogenation, and the sulfur impact on the MCH dehydrogenation reaction.

1. INTRODUCTION Hydrogen has been proposed as a clean energy carrier, helping to alleviate environmental pollution caused by fossil fuels. The future of this energy source is primarily dependent upon the development of suitable technology for safe and cost-effective storage and transport. Over the past few decades, different hydrogen storage alternatives have been intended, such as hydrogen compression in pressurized vessels, hydrogen liquefaction, hydrogen adsorption in metal hydrides, cryogenic storage with hydrogenadsorbing materials, and hydrogen storage in liquid organic hydrides (i.e., cycloalkane). The comparison between the different storage systems indicates that storing hydrogen in liquid organic hydrides has environmental, economical, technical, and social advantages.1
17 Environmentally, the system maintains a closed carbon cycle, which drastically reduces chemical and thermal pollutions caused by toxic gas emission. Economically, the system does not require heavy capital investment because the existing energy infrastructure is very compatible with it and recent economical analysis reveals a high feasibility.1
4 Technically, the previous implementation of this hydrogen storage system in mobile and stationary applications is promising, which encourages further research and development efforts.4
17 Socially, dealing with liquid organic hydrides is not that different from gasoline used today and more accepted publicly compared to gaseous and liquefied hydrogen. All of the previously mentioned reasons make the hydrogen storage in liquid organic hydrides a very competitive option for energy storage in the near future. The use of liquid organic hydrides in hydrogen storage and transportation has been initially investigated by the Euro-Quebec Hydrogen Project in the 1980s.18,19 The concept is based on reversible catalytic hydrogenation
dehydrogenation reactions. First, the hydrogen is chemically stored in an organic carrier r 2011 American Chemical Society

through a catalytic hydrogenation reaction. When the demand of energy exists, the hydrogen is extracted from the organic carrier by a catalytic dehydrogenation reaction and fed into fuel cells to generate electricity. Hydrogen storage in liquid organic hydrides has strongly influenced clean energy research in the past few decades because of its simple, safe, and feasible handling of hydrogen. The intensive research efforts have resulted in various liquid organic cycles for hydrogen storage. The most promising and famous liquid organic cycles are presented in the following section.

2. LIQUID ORGANIC HYDRIDE CYCLES A liquid organic hydride cycle consists of three components: naphthene, aromatic, and hydrogen. In the hydrogenation reaction, the aromatic compound is converted to naphthene, and in the dehydrogenation, the opposite should occur. For the cycle to be technically favorable and economically feasible, the hydrogenation
dehydrogenation reactions should demonstrate complete reaction reversibility, high selectivity, and high hydrogen density. The most famous liquid organic hydride cycles today are methylcyclohexane
toluene
hydrogen (MTH cycle), cyclohexane
benzene
hydrogen (CBH cycle), and decaline
naphthalene
hydrogen (DNH cycle). The hydrogen storage densities of these cycles are indicated in Table 1. 2.1. Methylcyclohexane
Toluene
Hydrogen (MTH) Cycle. The usage of methylcyclohexane (MCH) as a liquid organic carrier of hydrogen was first proposed by Sultan and Shaw in 1975.20 This concept of hydrogen storage is known today as the MTH cycle. Received: June 5, 2011 Revised: September 15, 2011 Published: September 16, 2011 4217

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Table 1. Gravimetric and Volumetric Contents of Liquid Organic Cycles1 gravimetric content

volumetric content

reaction

(wt %)

(kg H2/m3)

C7 H8 þ 3H2 S C7 H14

6.1

47.0

C6 H6 þ 3H2 S C6 H12

7.1

55.5

C10 H8 þ 5H2 S C10 H18

7.2

64.9

Many preceding studies4
12,21,22 considered the MTH cycle to be the most promising cycle for hydrogen storage because it is extremely reversible, highly selective, and free of carcinogenic products. The following sections will briefly illustrate the preceding efforts and the promising applications of the MTH cycle. 2.1.1. Stationary Application of the MTH Cycle. The economical and technical potential of the MTH cycle has encouraged the system implementation in mobile and stationary applications, such as hydrogen-powered vehicles and fuel-cell power stations, respectively. The concept of integrating the MTH cycle in a stationary application is based on balancing the mismatch in electricity supply and demand between seasons or time of day. For example, the surplus of summer electricity can be used in water electrolysis to generate hydrogen, which is then combined with toluene in an exothermic hydrogenation reaction to give MCH (eq 1). To match the high energy demand of the winter season, hydrogen will be recovered from the MCH by an endothermic dehydrogenation reaction (eq 2) and sent to a fuel cell power plant to generate electricity. C7 H8 þ 3H2 f C7 H14

ΔH250 °C ¼
214:1 kJ=mol

ð1Þ

C7 H14 f C7 H8 þ 3H2

ΔH450 °C ¼ þ 216:3 kJ=mol

ð2Þ

To make the application of the MTH cycle in the power station more efficient and feasible, the heat generated from fuel cells, during the re-electrification process, should be used in the endothermic dehydrogenation reaction, which normally occurs between 400 and 500 °C. Three re-electrification power plant systems were proposed for the integration with the MTH cycle. These re-electrification options are gas/steam turbine, low-temperature fuel cells, and high-temperature fuel cells.5,6 Scherer et al.5 have evaluated the three re-electrification alternatives to determine the best integrating option with the MTH cycle. Lowtemperature fuel cells (i.e., phosphoric acid fuel cell and polymer electrolyte fuel cell) were excluded because their operating temperature (200 °C) cannot provide adequate heat for the dehydrogenation of MCH. The gas/steam turbine option is capable of supplying adequate heat for the dehydrogenation reaction; however, the heat-transfer process is fairly complicated and expensive. A high-temperature fuel cell power plant is the best existing option that can supply the dehydrogenation reaction with the required energy. The two possible alternatives of high-temperature fuel cells are molten carbonate fuel cells (MCFC) and solid oxide fuel cells (SOFC). In these fuel cells, the chemical energy that is not converted to electrical power is converted to heat. The operating temperatures of the MCFC and SOFC are 650 and 950 °C, respectively. Efficiency analysis indicates that the integration of the MTH cycle with SOFC has a better process efficiency than MCFC because it provides more heat and does not require the extraction of CO2 from the outlet stream.

2.1.2. Mobile Application of the MTH Cycle. Hydrogen energy has attracted the automotive industry for many decades because it offers a clean and widely available source of energy. A serious research effort was initiated by Rudolph Erren in the 1930s when he started the development of hydrogen
gasoline dual-fuel vehicles. The interest in hydrogen as a source of energy rapidly diminished after World War II. Another wave of interest was initiated in the early 1970s by Japan, West Germany, and the United States. These efforts have resulted in many hydrogenfueled vehicles on which hydrogen was stored on board in the form of compressed gas, liquid hydrogen, or metal hydride.7 The compatibility of liquid organic hydrides with the existing transport-sector infrastructures (i.e., storage tanks, refueling stations, and transportation system) has given them economical, technical, and sociological advantages over the other hydrogen storage alternatives. In 1980, Taube and co-workers presented the notion of integrating the MTH cycle in hydrogen-powered vehicles, and in 1984, they successfully implemented the MTH cycle in a hydrogen-powered vehicle.8,9 According to their study, the concept of using a dehydrogenation catalytic reactor onboard a vehicle seems to be reasonable with respect to cost, size, weight, and maintenance. The effort of Taube et al.8,9 has established many technical and economical foundations for the integration of the MTH cycle in hydrogen-powered vehicles. Between 1985 and 1987, a second prototype of a hydrogenpowered vehicle was designed, constructed, and experimentally tested by Grunenfelder and Schucan.10 The geometrical size of the MTH dehydrogenation unit in the second prototype was reduced significantly, as compared to the one developed by Taube and co-worker. The promising results obtained from the second prototype truck have encouraged Grunenfelder and Schuncan to further improve the system and reduce the weight and size of the dehydrogenation unit by a factor of 3. Carrying out a highly endothermic reaction on-board a vehicle requires large amounts of heat. According to Taube et al.,8 the dehydrogenation heat equals one-third of the net combustion heat. To derive the catalytic reaction, Taube and co-workers suggested the extraction of the required heat from the exhaust gases (700 °C). However, the work carried out by Taube et al.9 indicated that the thermal energy from the exhaust gases did not provide sufficient heat to sustain the dehydrogenation reaction. To compensate for the lack of heat, Taube and co-workers proposed an auxiliary heating of the reactor that is provided from an external combustion of toluene. To reduce the toxic gas emission associated with toluene combustion, Kariya et al.11 suggested the oxidation of hydrogen to provide the thermal energy for the dehydrogenation reaction. Klvana et al.12 have proposed another approach for providing adequate heat to the dehydrogenation reactor by coupling the dehydrogenation reactor to the hydrogen engine. Klvana and co-workers studied three coupling orientations in laboratory scale. The first option is to couple each engine cylinder (total of six cylinders) with an annular catalytic bed, which is externally insulated to minimize heat dissipation. In an alternative orientation, the insulator is replaced by an exhaustgas heat exchanger to heat the exterior wall of dehydrogenation reactor. In a third coupling orientation, fins are employed to enhance the heat transfer from the engine cylinder to the exterior wall of the catalytic bed. Klvana et al.12 reported that the second and third orientations have shown a very promising effect on the dehydrogenation yield. The implementation of the MTH cycle in hydrogen-fueled vehicles constitutes an important contribution to environmental 4218

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Figure 2. Effect of hydrogen spillover and hydrogen recombination on the dehydrogenation of MCH over the Pt catalyst.

in the dehydrogenation reaction. Rahimpour et al.28 compared the dehydrogenation performance of decalin and cyclohexane in a thermally coupled reactor for Fischer
Tropsch synthesis (FTS) in gas-to-liquid (GTL) technology. Their results indicated that decalin-coupled configuration has a remarkably higher hydrogen production rate and hydrogen recovery yield than cyclohexanecoupled configuration. Other research groups have also investigated the dehydrogenation of decalin to naphthalene and reported high dehydrogenation activity.29
33 Figure 1. Schematic diagram for the implementation of the MTH cycle in mobile and stationary applications.

protection because it stops the emission of toxic gases, reduces the thermal pollution by a factor of 6 as compared to a gasolinedriven vehicle,8 and reduces the noise pollution in highly populated areas. The implementation of the MTH cycle in mobile and stationary applications is schematically illustrated in Figure 1. 2.2. Cyclohexane
Benzene
Hydrogen (CBH) Cycle. A number of research groups have investigated the CBH cycle.1,11,22
26 Cacciola et al.23 reported that the reaction reversibility and selectivity in the CBH cycle is higher than that of MTH and DNH cycles. Kariya et al.11,25 studied hydrogen production from cyclohexane, MCH, and decalin and reported that the highest production rate was obtained from the dehydrogenation of cyclohexane. Wang et al.22 compared the dehydrogenation of cyclohexane and MCH over Pt/Al2O3 and reported the turnover number (TON) of hydrogen production to be slightly higher in cyclohexane dehydrogenation (0.44 and 0.41 s
1, respectively). Despite the higher reversibility, better selectivity, and superior hydrogen storage density, researchers still prefer the MTH cycle over the CBH cycle because the latter contains benzene, which is carcinogenic. 2.3. Decalin
Naphthalene
Hydrogen (DNH) Cycle. Different research groups have studied the possibility of storing hydrogen energy in the DNH cycle. Cacciola et al.23 and Kariya et al.25 reported that the DNH cycle has poor reversibility. Kariya and co-workers also indicated that the DNH cycle is less selective compared to the CBH cycle. Another limitation of the DNH cycle is exemplified in the solid phase of naphthalene, which has a melting point of 218 °C. This constraint requires the system to be always heated to prevent any line blockage. Newson et al.,27 on the contrary, reported advancements in the catalytic process that allowed them to carry out the hydrogenation
dehydrogenation reactions of the DNH cycle interchangeably using one catalyst at different operating conditions. According to Newson and his colleagues, the single reactor engineering system and the advancement in catalytic reactions allowed them to replace the MTH cycle with the DNH cycle, which provides 38% extra hydrogen

3. KINETICS OF MCH DEHYDROGENATION MCH dehydrogenation to form toluene is an important reaction in the catalytic reforming process. Transition metals, such as Pt and Pd, are very active components with significant dehydrogenation ability.34 Product selectivity in dehydrogenation reactions is very crucial for industrial applications. The reaction selectivity is normally identified by the reaction mechanism, which is dependent upon the catalyst type. For instance, carrying out MCH dehydrogenation over Pt/NaY zeolite has illustrated low selectivity that yields different compounds, such as methylcyclohexene, methylcyclohexadiene, and toluene.35,36 In contrast, the dehydrogenation of MCH over Pt/Al2O3 has illustrated much higher selectivity toward toluene formation. Previous studies have shown that Pt/Al2O3 and its Pt
Re/Al2O3 successor are the best catalysts for MCH dehydrogenation in terms of activity, selectivity, and stability. However, the literature clearly demonstrates a considerable discrepancy with regard to the kinetic of MCH dehydrogenation over these two catalysts. This discrepancy is illustrated and reviewed in the following sections. 3.1. MCH Dehydrogenation over the Pt/Al2O3 Catalyst. The application of Pt catalysts in the dehydrogenation reaction was initially recognized by Zelinskii in 1911. The commercial application of Pt catalysts, however, did not start until 1949.37 Kariya et al.11 indicated that the good dehydrogenation ability of the Pt catalyst is mainly attributed to the rapid elimination of hydrogen from the reaction system, which shifts the chemical equilibrium in the favor of the products. Kariya and co-workers believe that the fast evolution of hydrogen in the Pt catalyst is credited to the effects of hydrogen spillover and hydrogen recombination, which are illustrated in Figure 2. The dehydrogenation of MCH over the Pt/Al2O3 catalyst has attracted many research groups over the past few decades. For example, Sinfelt et al.38 have investigated the kinetics of MCH dehydrogenation over Pt/Al2O3 and reported zero reaction order with respect to both MCH and hydrogen. Sinfelt and co-workers conducted the dehydrogenation reaction under the following conditions: temperature range of 315
372 °C, MCH partial 4219

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pressure of 0.07
2.2 atm, and hydrogen partial pressure of 1.1
4.1 atm. The zero reaction order with respect to MCH clearly suggests that the active platinum sites are completely covered by adsorbed MCH molecules, which makes the reaction rate independent of further MCH additions. The presence of toluene, on the other hand, had a small retarding effect on the dehydrogenation rate. Therefore, Sinfelt et al.38 suggested that the rate-determining step of MCH dehydrogenation is toluene desorption. On the basis of these observations, they derived the following rate expression for MCH dehydrogenation over Pt/Al2O3: rMCH f TOL ¼

kbPMCH 1 þ bPMCH

ð3Þ

The rate expression proposed by Sinfelt and co-workers37,38 clearly suggests that the reaction rate is approximately equal to the temperature-dependent variables at an elevated MCH partial pressure. These variables can be evaluated using eqs 4 and 5. The small inhibition effect of toluene reported by Sinfelt et al.37 has not been shown in the rate expression of eq 3. Hawthorn et al.39 applied the rate expression proposed by Sinfelt et al.37 within a mathematical model of a packedbed/heat-exchanger reactor for MCH dehydrogenation and reported a satisfactory result.

33000 11 k ¼ 4:2  10 exp ð4Þ RT b ¼ 2:1  10


10



30000 exp RT

ð5Þ

Wolf and Petersen40 have also derived a rate expression for MCH dehydrogenation over the Pt/Al2O3 catalyst with the assumption that the surface reaction is the rate-determining step. This rate expression is obtained from the Langmuir
Hinshelwood reaction mechanism and is illustrated in eq 6. The term “PS” in the adsorption group (denominator) stands for reversible poison, and it will be addressed later in the Kinetics of Catalyst Deactivation. rMCH f TOL ¼ ðkMCH f TOL KMCH PMCH Þ=ð1 þ KMCH PMCH þ KMCH KPS PMCH =PH þ KTOL PTOL Þ

ð6Þ

Jossens and Petersen41 presented very similar observations to those reported by Sinfelt and co-workers. According to them, the rate of MCH dehydrogenation appears to be first-order with respect to MCH partial pressure when the MCH concentration is below 6  10
7 gmol/cm3. At higher concentrations, however, the reaction order gradually decreases until it eventually approaches zero. Jossens and Petersen also reported that toluene formation inhibits the dehydrogenation reaction. According to them, an increase in the temperature strongly increases toluene inhibition, whereas an increase in hydrogen partial pressure significantly decreases toluene inhibition. In addition to Sinfelt et al.38 and Jossens and Petersen,41,42 toluene inhibition in MCH dehydrogenation over Pt/Al2O3 has been reported by many other researchers.43
45 Jossens and Petersen41 reported that toluene inhibition has an order of
1/3, while Andreev et al.43 reported that the inhibition order is
1/2. Jossens and Petersen42 and Jothimurugesan et al.44 believed that the retarding effect of toluene on the dehydrogenation rate is attributed to the competitive adsorption between MCH and toluene for the available active sites on the catalyst surface. Van Trimpont et al.46,47 have conducted a detailed kinetic study for MCH dehydrogenation over Pt/Al2O3. Van Trimpont

and co-workers examined the influence of MCH partial pressure on the dehydrogenation rate over the Pt/Al2O3 catalyst and reported three stages. Below 0.5 bar of MCH partial pressure, the concentration of MCH is relatively low and the dehydrogenation rate is linearly increasing with the MCH concentration, suggesting first-order kinetics. Between 0.5 and 1.5 bar, the dehydrogenation rate gradually levels off until the rate becomes almost constant (reaction order between 0 and 1). Above 1.5 bar, the dehydrogenation rate becomes independent of the MCH partial pressure; thus, the reaction order is zero. The results reported by Van Trimpont et al.46 extend those initially observed by Sinfelt et al.38 and Jossens and Petersen.41,42 Van Trimpont et al.46 also examined the effect of the hydrogen partial pressure on MCH dehydrogenation over the Pt/Al2O3 catalyst. They indicated that an increase in the hydrogen partial pressure sharply decreased the dehydrogenation rate on Pt/Al2O3. This inhibiting effect is inconsistent with the one reported by Sinfelt et al.,38 in which they reported insignificant impact of the hydrogen partial pressure on the MCH dehydrogenation rate over Pt/Al2O3. Van Trimpont and co-workers also researched the competitive adsorption between MCH, toluene, and n-heptane on the Pt/Al2O3 catalyst. Their assessment of competitive adsorption revealed that toluene does not compete with MCH for Pt active sites, which contradicts the previously mentioned observations. Van Trimpont and co-workers, however, observed competitive adsorption between n-heptane and MCH. Van Trimpont et al.46 derived several Hougen
Watson rate equations for MCH dehydrogenation based on different reaction mechanisms (Table 2). Most of these equations were rejected because their parameters have significant negative values. Only four reliable rate equations remained, and they are presented in Table 3 together with their corresponding rate-determining steps. The structures of the best four rate equations illustrate remarkable similarities. From the Hougen
Watson rate equation, Van Trimpont and co-workers concluded that the negative reaction order of hydrogen is not attributed to hydrogen competitive adsorption but to the involvement of surface intermediates in the rate-determining step. The adsorption terms of the rate equations in Table 3 do not involve toluene, which confirms the absence of competitive adsorption between MCH and toluene. Again, this finding contradicts the observations of Jossens and Petersen,42 Jothimurugesan et al.,44 and Pacheco and Petersen,45 in which they reported a retarding effect of toluene because of competitive adsorption. El-Sawi et al.48 studied the kinetics of MCH dehydrogenation over Pt/Al2O3 at atmospheric pressure and over the temperature range of 573
623 K. Their kinetic model is mainly based on a mechanistic model proposed by Jothimurugesan et al.44 for MCH dehydrogenation over Pt
Re/Al2O3. The estimation of the kinetic parameters for the monometallic catalyst was determined by applying nonlinear regression. The rate expression proposed by Jothimurugesan et al.44 and adapted by ElSawi and co-workers48 is shown in eq 7, and the rate-determining step is MCH adsorption. Although El-Sawi and co-workers obtained a good fit for their experimental results using this rate expression, the estimated activation energy happens to be lower than the one reported by Jothimurugesan et al.44 (17.92 kJ mol
1 compared to 51.9 kJ mol
1) and much lower than the one reported by Van Trimpont et al.46 (133.3 kJ mol
1). El-Sawi and co-workers attributed the low activation energy 4220

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value to the strong correlation between the pre-exponential factor and the activation energy. rMCH ¼ k

ðPMCH
PTOL PH 3 =Ke Þ ð1 þ KTOL PTOL Þ

ð7Þ

Chai and Kawakami49 proposed another kinetic model for MCH dehydrogenation over Pt/Al2O3. Their model is based on a single-site reaction mechanism (sequence III in Table 2)

and considered the dehydrogenation of methylcyclohexene to methylcyclohexadiene to be the rate-controlling step. The rate of dehydrogenation derived by Chai and Kawakami is given in eq 8. The rate expression obviously designates an improvement in the dehydrogenation rate as the MCH partial pressure increases and retardation in the rate as hydrogen partial pressure rises. The rate expression also suggests the absence of toluene inhibition. rMCH f TOL ¼ ðkKMCH1 KMCH PMCH =PH Þ=ð1 þ KMCH PMCH þ KMCH1 KMCH PMCH =PH þ KX KMCH1 KMCH PMCH =PH 2 Þ

Table 2. Reaction Sequence for the Dehydrogenation of MCH46 a sequence I: dual-site surface reactions, generation of atomic hydrogen MCH þ L T MCH-L

ð1Þ

MCH-L T A1-L þ H-L

ð2Þ

A1-L þ L T MCH1-L þ H-L

ð3Þ

MCH1-L þ L T A2-L þ H-L

ð4Þ

A2-L þ L T MCH2-L þ H-L

ð5Þ

MCH2-L þ L T A3-L þ H-L

ð6Þ

A3-L þ L T TOL-L þ H-L

ð7Þ

2H-L T H2 þ 2L

ð8Þ

TOL-L T TOL þ L

ð9Þ

sequence II: dual-site surface reactions, generation of molecular hydrogen MCH þ L T MCH-L

ð1Þ

MCH-L þ L T MCH1-L þ H2 -L

ð2Þ

MCH1-L þ L T MCH2-L þ H2 -L

ð3Þ

MCH2-L þ L T TOL-L þ H2 -L

ð4Þ

TOL-L T TOL þ L

ð5Þ

H2 -L T H2 þ L

ð6Þ

sequence III: single-site surface reactions, no adsorption
desorption of hydrogen MCH þ L T MCH-L

ð1Þ

MCH-L T MCH1-L þ H2

ð2Þ

MCH1-L T MCH2-L þ H2

ð3Þ

MCH2-L T TOL-L þ H2

ð4Þ

TOL-L T TOL þ L

ð5Þ

a

MCH, methylcyclohexane; MCH1, methylcyclohexene; MCH2, methylcyclohexadiene; TOL, toluene; Ai, intermediates; and L, active site.

ð8Þ Tsakaris50 investigated the dehydrogenation of MCH over a range of Pt catalysts at atmospheric pressure. A Langmuir
Hinshelwood and Hougen
Watson kinetic model was proposed, assuming the surface reaction as the rate-limiting step of the overall dehydrogenation reaction. The apparent activation energy is estimated to be 50 kJ mol
1 with a standard deviation of 6.02 kJ mol
1. Alhumaidan51 developed various power-law models for MCH dehydrogenation over Pt catalysts. The information obtained from the power-law models was extremely valuable and provided clues as to the form of function needed in developing more mechanistic and practically useful models. For example, the power-law models indicated that the reaction order with respect to MCH is approximately zero, while it is negative with respect to hydrogen. The power-law models also indicate that hydrogen inhibition increases with the total pressure. In addition to the previous observations, the power-law models show that the apparent activation energy of MCH dehydrogenation almost doubles as the pressure increases from 1 to 3 bar and then remains nearly constant as the pressure increases. Such observation has been attributed to a possible change in the reaction mechanism and the onset of structure sensitivity between the low- and high-pressure regions. Alhumaidan et al.52 used the observations of the power-law model to develop several mechanistic kinetic models for MCH dehydrogenation over Pt catalysts. The best fitting mechanistic model was the non-Langmuirian/noncompetitive Horiuti
Polanyi model. In this model, the Horiuti
Polanyi aromatic hydrogenation mechanism, which assumes an atomic hydrogen addition to aromatics on the catalyst surface, was applied in reverse to MCH dehydrogenation. The model also assumes that hydrogen and MCH molecules adsorb on two different types of sites to accommodate the observed near zero-order dependence of the reaction rate on MCH and the negative reaction

Table 3. Best Derived Hougen
Watson Rate Equations for MCH Dehydrogenation over the Pt/Al2O3 Catalyst46 rate-determining stepa

Hougen
Watson rate equation

a

rMCH f TOL ¼

kMCH f TOL ðPMCH
PTOL PH 3 =KMCH f TOL Þ=PH 3=2 ð1 þ KMCH PMCH þ KA2-L ðPMCH =PH 3=2 ÞÞ2

I(5)

rMCH f TOL ¼

kMCH f TOL ðPMCH
PTOL PH 3 =KMCH f TOL Þ=PH 5=2 ð1 þ KMCH PMCH þ KA3-L ðPMCH =PH 5=2 ÞÞ2

I(7)

rMCH f TOL ¼

kMCH f TOL ðPMCH
PTOL PH 3 =KMCH f TOL Þ=PH ð1 þ KMCH PMCH þ KMCH1 ðPMCH =PH ÞÞ2

II(3)

rMCH f TOL ¼

kMCH f TOL ðPMCH
PTOL PH 3 =KMCH f TOL Þ=PH 2 ð1 þ KMCH PMCH þ KMCH2 ðPMCH =PH 2 ÞÞ2

II(4)

Refer to Table 2. 4221

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Figure 3. Graphical comparison between fitted and observed conversions for the non-Langmuirian/noncompetitive Horiuti
Polanyi model.52

Table 4. Controversy about the Kinetics of MCH Dehydrogenation over Pt/Al2O3 rate-determining reference Sinfelt et al.

reaction rate

38 40

Wolf and Petersen

step

toluene inhibition

rMCH f TOL ¼ ðkbPMCH Þ=ð1 þ bPMCH Þ

toluene desorption

yes

rMCH f TOL ¼ ðkKMCH PMCH Þ=ð1 þ KMCH PMCH

surface reaction

yes

dehydrogenation of

no

þ KMCH KPS PMCH =PH þ KTOL PTOL Þ Van Trimpont et al.

46

rMCH f TOL ¼ ðkðPMCH
PTOL PH 3 =KMCH f TOL Þ=PH 2 Þ

methylcyclohexadiene

2

=ðð1 þ KMCH PMCH þ KMCH2 ðPMCH =PH ÞÞ Þ 2

to toluene

rMCH f TOL ¼ kðPMCH
PTOL PH 3 =Ke Þ=ð1 þ KTOL PTOL Þ

El-Sawi et al.48 49

Chai and Kawakami

rMCH f TOL ¼ ðkKMCH1 KMCH PMCH =PH Þ=ð1 þ KMCH PMCH þ KMCH1 KMCH PMCH =PH þ KX KMCH1 KMCH PMCH =PH Þ 2

MCH adsorption

yes

dehydrogenation of

no

methylcyclohexene to methylcyclohexadiene

rMCH f TOL ¼ ðkPMCH ð1
PTOL PH 3 =Ke PMCH ÞÞ

Tsakiris50

surface reaction

yes

dehydrogenation of MCH to

yes

2

=ðð1 þ KTOL PTOL þ KH PH Þ Þ 52

Alhumaidan et al.

ðrMCH f TOL Þo ¼ k=ð1 þ KH 0:5 PH 0:5 Þ

methylcyclohexene

order dependence upon hydrogen. To account for the increase in hydrogen inhibition with pressure, a non-Langmuirian adsorption isotherm was adopted, which assumes a nonlinear dependency between the adsorption equilibrium constant for hydrogen and the system pressure. The model also satisfactorily incorporates the reversible and irreversible deactivation kinetics. The statistical parameters and the graphical comparison between fitted and observed data (Figure 3) confirm that the model can successfully describe the MCH dehydrogenation reaction with a small number of well-determined and uncorrelated kinetic parameters. The previous discussion clearly demonstrates a considerable discrepancy in the literature with regard to the kinetics of MCH dehydrogenation over Pt/Al2O3. Examples of the conflicting and contradictory schemes that have been proposed in the literature are illustrated in Table 4.

3.2. Kinetics of MCH Dehydrogenation over the Pt
Re/ Al2O3 Catalyst. A great advancement was accomplished in 1968

when the platinum catalyst was promoted by the addition of rhenium (Re). The bimetallic reforming catalyst, Pt
Re/Al2O3, has shown higher stability and better selectivity toward aromatics than the monometallic Pt/Al2O3 catalyst.53,54 Furthermore, the addition of Re enhanced the dehydrogenation rate by promoting the cleavage of the C
H bond, improving the stability of the catalyst, and enhancing the desorption of hydrogen and aromatic products.11 The Pt
Re/Al2O3 catalyst normally contains equal amounts of Re and Pt (about 0.3 wt % of each). The significant results obtained from the Pt
Re catalyst have encouraged the development of many other bimetallic reforming catalysts, such as Pt
Sn/Al2O3, Pt
Ge/Al2O3, Pt
Rh/Al2O3, and Pt
Ir/Al2O3. However, the superior stability and selectivity of Pt
Re/Al2O3 4222

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Table 5. Best Derived Hougen
Watson Rate Equations for MCH Dehydrogenation over the Pt
Re/Al2O3 Catalyst46 Hougen
Watson rate equation

rate-determining step

rMCH f TOL ¼ ðkMCH f TOL ðPMCH
PTOL PH 3 =KMCH f TOL ÞÞ=ðð1 þ KMCH PMCH þ KTOL PTOL þ KnP7 PnP7 Þ2 Þ

I(2) and II(2)

rMCH f TOL ¼ ðkMCH f TOL ðPMCH
PTOL PH =KMCH f TOL ÞÞ=ð1 þ KMCH PMCH þ KTOL PTOL þ KnP7 PnP7 Þ

III(2)

3

led it to be the most widely used in reforming applications.55,56 Many researchers have tried to study the kinetics of MCH dehydrogenation over Pt
Re/Al2O3. For example, Jothimurugesan et al.44 applied the linear and nonlinear regression analysis to derive several Langmuir
Hinshelwood rate equations, which are based on both single- and dual-site reaction mechanisms. According to them, the best reaction mechanism starts with MCH adsorption followed by a series of consecutive reactions to yield methylcyclohexene, methylcyclohexadiene, and toluene (sequence III in Table 2). The rate equation of this reaction mechanism is given in eq 9. Jothimurugesan et al.44 reported that the adsorption of MCH is the rate-controlling step for the overall reaction kinetics and that the activation energy is 51.9 kJ mol
1. Jothimurugesan and co-workers studied the effect of product inhibition and reported that the presence of hydrogen did not affect the dehydrogenation rate, while the presence of toluene inhibited the dehydrogenation reaction. rMCH f TOL ¼

kðPMCH
PTOL PH 3 =Ke Þ 1 þ KTOL PTOL

ð9Þ

The kinetics of MCH dehydrogenation over Pt
Re/Al2O3 was also studied by Pacheco and Petersen.45 Their investigation resulted in a different rate expression than the one proposed by Jothimurugesan et al.44 (eq 10). However, they still agree that MCH adsorption is the rate-determining step and the activation energy is around 58.6 kJ mol
1. Pacheco and Petersen45 investigated toluene inhibition on MCH dehydrogenation over Pt
Re/Al2O3 and reported no inhibition effect, which again contradicts the finding of Jothimurugesan and co-workers.44 rMCH f TOL ¼

ko PMCH PMCH 1 þ KTOL 1=2 PH

ð10Þ

Pal et al.57 derived various rate expressions for the dehydrogenation of MCH over Pt
Re/Al2O3. According to Pal and coworkers, the best rate expression derived is similar to the one proposed by Jothimurugesan et al.44 This dehydrogenation model is based on a single-site mechanism, and its rate expression is given in eq 11. In agreement with Jothimurugesan et al.44 and Pacheco and Petersen,45 Pal et al.57 reported the adsorption of MCH to be the rate-controlling step for MCH dehydrogenation over Pt
Re/Al2O3. The activation energy reported by Pal and his colleagues is 56.4 kJ mol
1, which is very close to the values reported by Jothimurugesan et al.44 and Pacheco and Petersen45 but lower than the value proposed by Jossens and Petersen42 (71.1 kJ/mol). Pal and co-workers also indicated that the addition of 10% toluene to MCH had a retarding effect on the rate of dehydrogenation, especially when hydrogen was the carrier gas. Conversely, when nitrogen was the carrier gas, toluene inhibition was negligible. rMCH f TOL ¼

kðPMCH
PTOL PH 3 =Ke Þ 1 þ KTOL PTOL

ð11Þ

Van Trimpont et al.46 also researched the kinetics of MCH dehydrogenation over the Pt
Re/Al2O3 catalyst. The comparison between the performance of Pt/Al2O3 and Pt
Re/Al2O3 indicated striking similarities and important differences. For example, the influence of the MCH partial pressure on the dehydrogenation over the Pt
Re/Al2O3 catalyst was very similar to that of Pt/Al2O3. At low MCH partial pressure (below 0.5 bar), the reaction rate linearly increased with the MCH pressure, suggesting first-order kinetics. In the range between 0.5 and 1.5 bar of MCH, the reaction rate gradually tended toward an asymptotic value. Above 1.5 bar of MCH pressure, the reaction rate lay on the asymptote, becoming independent of the MCH partial pressure and suggesting a rate that becomes zero-order with respect to MCH. In contrast to Pt/Al2O3, the influence of the hydrogen partial pressure on the dehydrogenation rate was negligible over Pt
Re/Al2O3. This contrasting behavior between the mono- and bimetallic catalysts was attributed to the partial substitution of platinum by rhenium, which shifted the rate-determining step from the dehydrogenation of methylcyclohexadiene into toluene to the dehydrogenation of MCH into methylcyclohexene. According to Van Trimpont et al.,46 the adsorption of methylcyclohexadiene is inversely related to the hydrogen partial pressure, whereas the adsorption of MCH is independent of the hydrogen partial pressure. Pacheco and Petersen45 reported that the apparent order of reaction with respect to hydrogen changes from 0 to 1/2 when the MCH reaction order approaches zero. Van Trimpont et al.46 also investigated toluene inhibition on MCH dehydrogenation over Pt
Re/Al2O3 and reported a significant increase in competitive adsorption between toluene and MCH over Pt active sites, as compared to the monometallic Pt catalyst. Van Trimpont and co-workers46 investigated the kinetics of MCH dehydrogenation over Pt
Re/Al2O3 and derived the corresponding Hougen
Watson rate equations using the previously proposed reaction mechanisms (Table 2). The most reliable rate equations for MCH dehydrogenation over Pt
Re/Al2O3, according to Van Trimpont and co-workers, are illustrated in Table 5. Reaction mechanisms I and II produced identical rate equations, while mechanism III only differs in the exponent of the adsorption term. The rate-determining step for all mechanisms was the dehydrogenation of MCH into methylcyclohexene. The n-heptane coefficient appears in the adsorption term because n-heptane was added intentionally to the feed (MCH) to evaluate the competitive adsorption. To summarize Van Trimpont’s observations, a brief comparison between the dehydrogenation of MCH over Pt/Al2O3 and Pt
Re/Al2O3 catalysts is given in Table 6. Chai and Kawakami49 reported that their MCH dehydrogenation model, previously proposed for Pt/Al2O3, is also valid for Pt
Re/Al2O3 catalysts. In their kinetic model, they assumed that the dehydrogenation of methylcyclohexene to methylcyclohexadiene is the rate-controlling step for both catalysts. Alhumaidan et al.52 also reported that the non-Langmuirian/noncompetitive Horiuti
Polanyi model, proposed for Pt/Al2O3, has successfully fit the data of Pt
Re/Al2O3 and Pt
Pd/Al2O3 catalysts. 4223

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Table 6. Comparison between the Dehydrogenation of MCH over Pt/Al2O3 and Pt
Re/Al2O3 Catalysts Based on Van Trimpont et al.46 Findings comparison aspect

Pt/Al2O3 catalyst

Pt
Re/Al2O3 catalyst

(n)MCH = 1 (PMCH < 0.5 bar)

(n)MCH = 1 (PMCH < 0.5 bar)

reaction order with respect to H2

(n)MCH = 0 (PMCH > 1.5 bar) inhibits the reaction

(n)MCH = 0 (PMCH > 1.5 bar) no inhibition effect

toluene inhibition

no inhibition effect

inhibits the reaction

activation energy

133.3 kJ mol
1

195.8 kJ mol
1

reaction mechanism

dual site is favored

single site is favored

rate-controlling step

dehydrogenation of methylcyclohexadiene into toluene

dehydrogenation of MCH into methylcyclohexene

reaction order with respect to MCH

Table 7. Controversy about the Kinetics of MCH Dehydrogenation over Pt
Re/Al2O3 reference

reaction rate

rate-determining step

toluene inhibition

MCH adsorption

yes

Pacheco and Petersen45

rMCH f TOL ¼ ðkðPMCH
PTOL PH =Ke ÞÞ=ð1 þ KTOL PTOL Þ
 PMCH rMCH f TOL ¼ ðko PMCH Þ= 1 þ K2 1=2 PH

MCH adsorption

no

Pal et al.57

rMCH f TOL ¼ ðkðPMCH
PTOL PH 3 =Ke ÞÞ=ð1 þ KTOL PTOL Þ

MCH adsorption

yes

rMCH f TOL ¼ ðkMCH f TOL ðPMCH
PTOL PH =KMCH f TOL ÞÞ

dehydrogenation of MCH to methylcyclohexene

yes

dehydrogenation of MCH to methylcyclohexene

yes

dehydrogenation of methylcyclohexene to

no

Jothimurugesan et al.

44

Van Trimpont et al.

46

Van Trimpont et al.

46

3

3

=ð1 þ KMCH PMCH þ KTOL PTOL þ KnP7 PnP7 Þ rMCH f TOL ¼ ðkMCH f TOL ðPMCH
PTOL PH 3 =KMCH f TOL ÞÞ 2

=ðð1 þ KMCH PMCH þ KTOL PTOL þ KnP7 PnP7 Þ Þ 49

Chai and Kawakami

rMCH f TOL ¼ ðkKMCH1 KMCH PMCH =PH Þ=ð1 þ KMCH PMCH þ KMCH1 KMCH PMCH =PH þ KX KMCH1 KMCH PMCH =PH 2 Þ

methylcyclohexadiene Alhumaidan et al.52

ðrMCH f TOL Þo ¼ k=ð1 þ KH 0:5 PH 0:5 Þ

dehydrogenation of MCH

yes

to methylcyclohexene

As in the case of the Pt/Al2O3 catalyst, the literature has illustrated a disagreement among researchers about the kinetics of MCH dehydrogenation over Pt
Re/Al2O3. This discrepancy is summarized in Table 7.

4. KINETICS OF CATALYST DEACTIVATION Catalyst deactivation occurs by different mechanisms, such as fouling, impurity poisoning, and sintering. Developing a catalyst that can maintain high stability requires a deep understanding of the deactivation phenomena. Catalyst deactivation can be either studied from a fundamental or an empirical approach.58 The fundamental approach presumes that the deactivation depends upon kinetic variables, such as temperature, pressure, liquid hourly space velocity (LHSV), and species concentrations. The empirical approach, on the other hand, assumes that the deactivation is a function of time on stream or fouling on the catalyst. The deactivation function is normally represented as a simple multiplier in front of the rate of reaction. 4.1. Deactivation of Pt/Al2O3 during MCH Dehydrogenation. The deactivation kinetics of Pt/Al2O3 catalysts during

MCH dehydrogenation was investigated by many researchers over the past few years. Wolf and Petersen40 studied the deactivation mechanism and the kinetics of poisoning and reported that the poisoning rate is inversely proportional to the hydrogen partial pressure. Wolf and Petersen also suggested two types of poison structures: one reversible and the other irreversible. The relative amount of each poison structure depends upon the hydrogen partial pressure and the time on stream. At a low partial

Figure 4. Self-poisoning mechanisms.

pressure of hydrogen, the irreversible poison structure is more likely to form. According to Wolf and Petersen, the dependence of the poisoning rate upon the hydrogen partial pressure suggests a self-poisoning mechanism, instead of an impurity-poisoning mechanism. Catalyst deactivation by self-poisoning may occur by parallel, series, or parallel
series (triangular) mechanisms, as illustrated in Figure 4. On the basis of these poisoning mechanisms, Wolf and Petersen40 reported two phases of deactivation in MCH dehydrogenation over Pt/Al2O3. The initial phase of deactivation is due to parallel self-poisoning and is characterized by a rapid rate of deactivation. The second phase of deactivation, on the other hand, is caused by uniform self-poisoning and has a slower rate of deactivation. The deactivation function of Wolf and Petersen is presented as a simple multiplier (the curly braces) in front of the rate expression of eq 12. Sto, P, and W are poisoning kinetic terms that represent the initial site concentration, the concentration of 4224

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Figure 5. Six-site fouling mechanism as a parallel pathway to the main reaction mechanism. An asterisk (/) represents an active Pt site.58

sites occupied by reversible poison, and the concentration of sites occupied by the irreversible poison, respectively. rMCH f TOL ¼ ðf½Sto
ð½P þ ½WÞgkMCH f TOL KMCH PMCH Þ= ð1 þ KMCH PMCH þ KMCH KPS PMCH =PH þ KTOL PTOL Þ

ð12Þ

The deactivation of Pt/Al2O3 in the dehydrogenation of MCH was also investigated by Jossens and Petersen.41 In agreement with Wolf and Petersen,40 they observed two phases of deactivation. The initial phase of deactivation is rapid, with a time period of minutes. In addition, it is characterized by low activation energy (33.5 kJ mol
1) and totally reversible deactivation in a stream of hydrogen. The rate of deactivation of the initial phase was nearly zero-order with respect to MCH and toluene partial pressures. The second phase of deactivation, on the contrary, is characterized by a slower rate of deactivation, a high activation energy (163.3 kJ mol
1), and partially reversible deactivation in a stream of hydrogen. The rate of deactivation in the second phase is proportional to the MCH dehydrogenation rate. Pacheco and Petersen58 proposed a deactivation mechanism and an empirical fouling correlation for the dehydrogenation of MCH. The fouling mechanism starts with a sequential adsorption of MCH on a group of six active sites (a sextet of sites). The series of dissociative adsorptions causes the MCH molecule to lose all 11 hydrogen atoms in its ring to form a multiply bound surface carbon skeleton. The carbon skeleton reacts with a gas-phase toluene molecule, in a Rideal
Eley manner, to form a fouled sextet, as illustrated in Figure 5. The rate-determining step for this fouling mechanism is the surface reaction between toluene and the multiply bound carbon skeleton. The apparent activation energy for the fouling reaction is 309.3 kJ mol
1. Pacheco and Petersen58 proposed a hyperbolic deactivation function for the six-site fouling mechanism a ¼ ð1 þ ktÞ
1=5 k¼

30kf K6 PMCH ðKTOL Þ6 PTOL PH 11=2

ð13Þ ð14Þ

where kf is the rate constant for the rate-determining step of the fouling reaction and K6 is the adsorption equilibrium constant for the fouling precursor. This hyperbolic deactivation function suggests that the fouling rate is directly related to the MCH partial pressure and inversely related to the hydrogen partial pressure. The deactivation function also suggests the inhibition of fouling by toluene. The deactivation rate proposed by Pacheco and Petersen58 fits the deactivation data quite well at a high level of catalyst activity. However, as the catalyst activity declines, the curve of experimental data becomes steeper and deviates from the predicted deactivation. This behavior suggests that a successful fouling model should not be based on a single reaction order but on a variable reaction order to represent the fouling data over a wider range of catalyst activities. The variable reaction order implies a succession of parallel fouling reactions that are competing to deactivate the catalyst. The dominant fouling mechanism, among the competing fouling reactions, changes with activity. Pacheco and Petersen59 modified their previous sextet (six-site) fouling mechanism to fit the fouling data below an activity level of 40%. Their new fouling model is based on a multiplet fouling mechanism, proposed initially by Herington and Rideal in 1944. To explain the discrepancy between fouling data and the sextet model at the lower range of activity, the modified fouling mechanism suggests that the fouling occurs on smaller multiplets as the sextets become exhausted. The reaction sequence of the multiplet fouling mechanism is similar to the sextet mechanism, except that each intermediate species is treated as an independent source of fouling. On a fresh catalyst, the initial deactivation mainly occurs on a six-site multiplet (sextet) until the surface is depleted of that multiplet. The fouling mechanism then proceeds toward the five-site multiplet, and the process continues until it reaches the two-site multiplet (doublet fouling). The fouling reaction prefers the higher order multiplets because they have a lower fouling activation energy. According to Pacheco and Petersen, the initial rate of the sextet fouling is 1000 times faster than that of the doublet fouling. A summary of the multiplet fouling mechanism is indicated in Figure 6. Chai and Kawakami49 proposed a different mechanism for the deactivation of Pt/Al2O3 in MCH dehydrogenation. They suggested that methylcyclohexadiene is adsorbed in a different way from that in the main reaction to form a coke precursor (X). 4225

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reversible by overnight reduction under hydrogen, while the cumulative deactivation was mostly irreversible by reduction. The shortterm deactivation has been empirically determined by introducing two time scales into the apparent rate constant to account for the daily and cumulative deactivation k ¼ ko expð
kd td Þ expð
kc tc Þ

Figure 6. Multiplet fouling mechanism.59

The polymerization of “n” adjacently adsorbed precursors is assumed to be the rate-controlling step of the deactivation reaction. According to Chai and Kawakami, parameter “n” can be interpreted in different ways, such as the number of adjacently adsorbed coke precursor molecules, the number of active sites involved in the coke formation, or the order of the deactivation reaction. Chai and Kawakami49 reported that the value of “n” falls between 2.6 and 2.8 in a Pt/Al2O3 catalyst. Corma et al.35 also reported that “n” is equal to 3 in the dehydrogenation of MCH over the Pt
NaY catalyst. The deactivation model proposed by Chai and Kawakami is summarized in eqs 15
17. a ¼ ð1 þ KtÞ1
n

ð15Þ

K ¼ nðn
1Þkd ðKX KMCH1 KMCH PMCH =ðPH 2 λÞÞn

ð16Þ

λ ¼ 1 þ KMCH PMCH þ KMCH1 KMCH PMCH =PH þ KX KMCH1 KMCH PMCH =PH 2

where kd and kc are the rate constants for daily and cumulative deactivation, respectively. The daily deactivation rate constant (kd) was obtained by incorporating a time-dependent factor into a simple power-law model. It was found that kd is a strongly decreasing function of the total pressure (P), although the particular relationship is catalyst-dependent. The rate constant of cumulative or irreversible deactivation, on the other hand, was determined by fitting the kinetic data to the non-Langmuirian/ noncompetitive Horiuti
Polanyi model, which indicates negligible irreversible deactivation (kc = 0) for the majority of catalysts. Alhumaidan et al.60 also studied the long-term deactivation, associated with life or deactivation tests, of the Pt catalyst in MCH dehydrogenation. Alhumaidan and co-workers successfully include the long-term deactivation in the non-Langmuirian/ noncompetitive Horiuti
Polanyi model. This has been successfully achieved by assuming that the long-term deactivation is the sum of two exponential decay processes, one rapid and completed in a few days and the other slow, with a half-life of weeks to months (eq 19). In this equation, “X” represents the observed conversion, “t” is the time-on-stream, and “a, b, c, and d” are empirical constants that exemplify the following: a, the population of the rapidly deactivated sites; b, the first-order decay constant of the rapidly deactivated sites (h
1); c, the population of the slowly deactivated sites; and d, the first-order decay constant of the slowly deactivated sites (h
1). Figure 7 illustrates the deactivation data fitting of various Pt catalysts by the doubleexponential decay model. In their study, Alhumaidan and coworkers indicated that an optimization of catalyst stability can be achieved by reducing the duration and sharpness of the fast deactivation phase, which allows the catalyst to operate at higher activity in the slow deactivation phase.

ð17Þ

In agreement with Chai and Kawakami,49 García de la Banda et al.36 believed that the precursor to coke formation is not toluene but the partially dehydrogenated products, such as methylcyclohexene and methylcyclohexadiene. García de la Banda and co-workers consolidate their argument by indicating that the deactivation in methylcyclohexene dehydrogenation is 6 times higher than the deactivation in MCH dehydrogenation. These observations are in agreement with the hypothesis that assigned unsaturated molecules (i.e., olefins) are the precursors of coke formation.41 Alhumaidan et al.52 has successfully described the kinetic of MCH dehydrogenation over Pt catalysts using the non-Langmuirian/noncompetitive Horiuti
Polanyi model. In their model, they have successfully included the short-term deactivation, which accounts for both the daily and cumulative deactivation observed in the catalyst activity tests. The daily deactivation represents the extent of drop in activity over a 1 day experiment, while the cumulative deactivation is the sum of all irreversible deactivations of daily experiments. The extent of the loss of conversion for the daily deactivation was observed by a repeat run at the end of each experiment, while the extent of cumulative deactivation was determined by carrying out repeat experiments over longer time intervals. The daily deactivation was largely

ð18Þ

X ¼ a expð
btÞ þ c expð
dtÞ

ð19Þ

4.2. Deactivation of Pt
Re/Al2O3 during MCH Dehydrogenation. An impressive extension in the life of the monome-

tallic Pt/Al2O3 catalyst was achieved by adding Re to form a bimetallic reforming catalyst. Over the past few decades, scientists proposed different rational explanations to justify this fundamental catalysis phenomenon that has significantly advanced many conversion processes. The analogies and contrasts of various studies are illustrated below to signify the diverse understandings of fouling over Pt
Re/Al2O3 in MCH dehydrogenation. Engels and Lehmann61 and Leprince62 reported that Re addition to the Pt monometallic catalyst inhibits the loss of the active surface area by reducing Pt sintering. Bertolacini and Pellet,63 on the other hand, believed that Re addition reduces the rate of coke formation by converting coke precursor molecules to harmless compounds. Sachtler64 disagreed with the hypothesis proposed by Bertolacini and Pellet and suggested that Re addition does not change the rate of coke formation but the nature of this coke. Sachtler believed that the association of Pt with Re modifies the nature of carbonaceous residues by impeding the formation of graphitic structures. According to Somorjai and Blakeley,65 4226

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Figure 7. Fitting the deactivation data of tested Pt catalysts by the double-exponential decay model.60

inhibiting the reorganization of carbonaceous fragments to form a graphitic structure depends upon the topography of the active metal surface. They reported that flat crystal faces of platinum are easily covered by coke, as compared to the stepped crystal faces. Flat crystal faces are more likely to exist in large platinum particles than smaller particles. Doolittle et al.54 also reported that the rate of deactivation increases as the size of platinum ensembles increases. Coughlin et al.66,67 introduced a model that explains the surface reaction and the deactivation phenomena of Pt and Pt
Re catalysts during MCH dehydrogenation. To differentiate between the effect of metallic and acidic functions, they neutralized the acidic sites on a γ-Al2O3 support (with a solution of 0.05% NaOH) to reduce the cracking activity. Nevertheless, the results still indicated a substantial cracking activity. Coughlin and co-workers thereby concluded that not only is the support acidity responsible for the cracking reactions but the metallic function also plays an important role in hydrogenolysis activity (the rupture of heteroatom C bonds by the addition of hydrogen). Coughlin and co-workers explained this using the concept of kink (edge) and terrace (surface) sites of the Pt crystals (or ensembles). They associated the hydrogenolysis activity with the kink sites and the dehydrogenation activity with the terrace sites of the Pt ensembles. According to them, the metal crystal is initially activated by the hydrogenolysis reaction at the kink sites.

Further hydrocracking reactions form carbonaceous layers that deactivate the kink sites to a significant extent. These carbonaceous layers extend gradually to cover the terrace sites and then rearrange to form graphitic deposits. The presence of Re in the kink and terrace sites prevents the rapid deactivation by directing the hydrogenolysis selectivity toward gaseous products rather than carbonaceous deposits.66 Doolittle et al.54 also reported that the strong Re
C bond enhances the breakage of C
C bonds and prevents the formation of graphitic deposits. Pacheco and Petersen59 tried to explain the drop in the fouling rate in Pt
Re/Al2O3 by referring to the multiplet fouling mechanism. They believed that the addition of Re drastically reduces the number of multiplets (responsible for fouling as previously indicated) and increases the single-site population. Biloen et al.,68 Van Trimpont et al.,47 and Chai and Kawakami49 also believed that the addition of Re divides the large Pt ensembles into smaller ones, which reduces the possibility of having dual or triple (2.7 on average) adjacent Pt sites occupied by a coke precursor molecule. This brief literature discussion indicates that the majority of deactivation models have attributed the stability enhancement in the bimetallic catalyst to the partitioning of Pt particles by Re. Jossens and Petersen42 have compared the deactivation rates of Pt/ Al2O3 and Pt
Re/Al2O3 catalysts during MCH dehydrogenation. 4227

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Figure 8. Fitting the deactivation data of tested Pt
Re catalysts by the double-exponential decay model.60

They reported that the addition of Re has not affected the initial phase of deactivation but has significantly reduced the deactivation rate in the long-term deactivation period (second phase). Pacheco and Petersen45 compared the fouling kinetics in MCH dehydrogenation over Pt/Al2O3 and sulfided Pt
Re/Al2O3. According to them, the specific fouling rates kf (min
1) of the mono- and bimetallic catalysts are given in eqs 20 and 21, respectively. The comparison between the specific fouling rates indicates a different toluene dependence. The toluene dependence has a negative first-order effect on the monometallic catalyst, while it has a positive first-order effect on the sulfided bimetallic catalyst. This finding indicates that the fouling reaction in Pt/Al2O3 is inhibited by toluene formation. This is consistent with the main reaction over Pt/Al2O3, which is inhibited by toluene as well. According to Pacheco and Petersen,45 these observations suggest that the extent of surface coverage by toluene is significantly different in the two catalysts. In other words, the competitive adsorption between toluene and MCH is more noticeable in the monometallic catalyst than the bimetallic catalyst. Pacheco and Petersen attributed the reduction of toluene inhibition in sulfided Pt
Re/Al2O3 to the effective partitioning of Pt sites by Re and sulfur, which caused a reduction in surface coverage.   PMCH
37000 kf ¼ 1:3  1040 exp ð20Þ T PTOL PH 5:5 kf ¼ 6:8  1026

  PMCH PTOL
39000 exp T PH 1:3

ð21Þ

Pal et al.57 further investigated the deactivation of Pt
Re/Al2O3 during MCH dehydrogenation. They proposed a deactivation model that is based on a dual-site mechanism with the assumption that the formation of a coke precursor is the rate-controlling step. Their model suggests that the coke precursor is formed from two adjacently adsorbed MCH molecules. The model differentiated between the adsorption of MCH in coke formation and the adsorption of MCH in the dehydrogenation reaction. Pacheco and Petersen59 also suggested that a coke precursor is formed from two adjacently adsorbed MCH molecules; however, they believe the adsorption of MCH to be similar in both dehydrogenation and deactivation reactions. Pal et al.57 determined the activation energy for the deactivation reaction to be 146.6 kJ mol
1, which is close to the value obtained by Jossens and Petersen42 (163.2 kJ/mol).

The activation energy of the deactivation reaction is much higher than that of the dehydrogenation reaction (56.4 kJ mol
1), which suggests that the deactivation reaction is much more sensitive to temperature and has a slower rate, as compared to the dehydrogenation reaction. The deactivation model of Pal et al.57 is summarized in eqs 22
24, where kd and KMCH* are the rate constant of the deactivation reaction (h
1) and the equilibrium adsorption constant for MCH to yield the coke precursor (atm
1), respectively. Pal and co-workers also reported that the effect of toluene on the deactivation reaction is insignificant. This finding confirmed that the coke precursor is MCH and not toluene (deactivation reaction is parallel to the main reaction).


da ¼ ððkd ðKMCH Þ2 PMCH 2 Þ=ð1 þ KTOL PTOL þ KMCH PMCH Þ2 Þa2 dt

ð22Þ kd ¼ 5:64  1011 expð
17500=TÞ

ð23Þ

KMCH ¼ 0:014 expð3625=TÞ

ð24Þ

60

Alhumaidan et al. applied the double-exponential decay model, proposed for Pt catalysts, to empirically fit the long-term deactivation of Pt
Re catalysts in MCH dehydrogenation. Alhumaidan and co-workers reported a good fit, as illustrated in Figure 8.

5. EXCEEDING THERMODYNAMIC EQUILIBRIUM IN MCH DEHYDROGENATION The reversible dehydrogenation reaction of naphthenes to aromatics (i.e., MCH to toluene) is strongly limited by thermodynamic equilibrium. To enhance the system efficiency, one of the reaction products must be removed to shift the equilibrium condition. This shift in the equilibrium condition enhances the equilibrium conversion and decreases the severity of operating conditions. Ali et al.69 have attempted to exceed the equilibrium limitation of the MCH dehydrogenation reaction using a catalytic membrane reactor (shell-tube configuration). The membrane reactor contained a sulfided Pt/Al2O3 catalyst in the shell and a tubular Pd77Ag23 membrane (0.1 mm thick) sealed in the center. The use of the membrane allowed for selective in situ removal of hydrogen. The permeation of hydrogen through the membrane reduced the thermodynamic constraint and significantly enhanced the toluene and hydrogen yields to an extent that exceeded the equilibrium 4228

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yield (reaction limitation). A comparison between membrane and non-membrane reactors (conventional packed-bed reactor) indicated a much higher conversion for the membrane reactor that surpassed equilibrium conversion 4 times. Furthermore, the high conversion rate of the membrane reactor was achieved at a relatively low operating temperature, as compared to the non-membrane reactor. The experimental work of Ali et al.69 was conducted at temperatures of 573
673 K, pressures of 5
20 bar, and LHSV of 2
12 h
1. To operate at higher pressure and temperature, Ali and Baiker70 suggested the use of a thicker wall membrane. The work of Ali et al.69 was conducted with a catalyst diluted 10 times with inert particles. Catalyst dilution was intended to reduce the endothermic effect and maintain an isothermal operation, which is rather simple to model. Ali and Baiker71 derived another kinetic rate equation for MCH dehydrogenation using experimental data from a non-isothermal (undiluted catalyst-bed) membrane reactor. This rate equation was employed in a membrane reactor mathematical model. The simulation results of toluene and hydrogen yields fitted well to the experimental values. The MCH dehydrogenation rate equation proposed by Ali and Baiker71 is based on a rate equation developed by Manser in 1992 (eq 25). To best-fit the non-isothermal experimental data from the membrane reactor, the pre-exponential factor [kmol s
1 (kg of catalyst)
1 kPa
1] and the activation energy (J mol
1) of the rate constant (k) were re-evaluated. The re-evaluated rate constant of the membrane reactor is given in eq 26. " # PTOL ðPH2 Þ3 ð25Þ rMCH f TOL ¼ kPMCH 1
Keq PMCH k ¼ 9:0493  10


8



123520 1 1
exp R T 633:15

ð26Þ

Ali and Baiker71 calculated the equilibrium constant of eq 25 from the Akyurtlu and Stewart correlation, which is illustrated in eq 27. The first term in eq 27 represents the temperaturedependent equilibrium constant at 650 K. According to Schildhauer et al.,72 the correlation developed by Akyurtlu and Stewart revealed a lack of fit to experimental data. Schildhauer and coworkers proposed a new correlation for the equilibrium constant, to resolve the lack of fit, and is given in eq 28. 

Keq ¼ ð4:61  109 kPa3 Þexp


 ð
216350 J=molÞ 1 1
R T 650 K

ð27Þ 

Keq


 ð
217650 J=molÞ 1 1 ¼ ð3:60  10 kPa Þexp
R T 650 K 9

3

ð28Þ The concept of using a H2-selective membrane reactor in MCH dehydrogenation has attracted more attention because it provides a higher conversion rate, greater product yields, and simpler product separation. Ferreira-Aparicio et al.73 also studied the performance of a membrane reactor for the dehydrogenation of MCH. Three fundamental points were considered in their study: the membrane synthesis and permeance, the effect of operating conditions on the conversion rate and product yields, and the performance of the Pt/Al2O3 catalyst. To selectively separate the hydrogen product, Ferreira-Aparicio and co-workers used a

palladium/stainless-steel membrane with high permeance values and high thermal and mechanical resistance. To operate at optimum conditions, the effect of several operating parameters on the conversion level was explored, such as the MCH flow rate, the reaction temperature, the system pressure, and the sweep gas flow rate. According to Ferreira-Aparicio et al.,73 the temperature range of 573
623 K ensures a high conversion rate, a yield close to 100%, and a very limited deactivation of the catalyst. In a different research effort, Ferreira-Aparicio et al.74 studied the applicability of using a porous vycor glass membrane to enhance the dehydrogenation of MCH to toluene. The study indicated that the porous vycor membrane has lower hydrogen selectivity, as compared to a palladium-based membrane. However, considering the fact that a palladium membrane requires a high operating temperature (573 K) to maintain the high selectivity, the porous vycor membrane would be a good alternative for processes operating at relatively lower temperature. In addition to the use of membranes in shifting the equilibrium condition toward the dehydrogenation reaction, other research groups tried to exceed the dehydrogenation equilibrium using the superheated liquid-film-type catalysis. For example, Shinohara et al.,29 Hodoshima et al.,30
32 and Sebastian et al.33 investigated the dehydrogenation of decalin to naphthalene under liquid film conditions, using carbon-supported platinum-based catalysts (Pt/C), and reported an efficient catalytic dehydrogenation. Hodoshima et al.30
32 compared the dehydrogenation reaction under both a liquid film state and a suspended state and confirmed that the catalytic dehydrogenation under reactive distillation and a liquid film state easily surpasses the equilibrium conversion at a relatively low dehydrogenation temperature (200
280 °C). Hodoshima and co-workers attributed the enhancement in the dehydrogenation activity under reactive distillation conditions and a liquid film state to the superheated catalyst layer, which accelerates the dehydrogenation rate and the desorption of product naphthalene. Hodoshima et al.30 indicated that the highest conversion is achieved when the reactant/catalyst ratio is 4 mL g
1, while Sebastian et al.33 suggested 2.7 mL g
1. Kariya et al.11 also observed highly efficient evolution of hydrogen from cycloalkanes over the Pt/C catalyst under wet
dry multiphase conditions. Kariya and co-workers attributed the high dehydrogenation rate under wet
dry multiphase conditions to the feeding efficiency of the reactant, the high catalyst surface temperature, and the efficient removal of products (i.e., hydrogen and aromatics) from the reaction system. Kariya et al.11 indicated that the optimum state of wet
dry multiphase conditions can be obtained by careful control of the reaction temperature, reactant, catalyst support, and reactant/catalyst ratio. According to Kariya and coworkers, the highest rate of hydrogen formation is obtained when the reactant/catalyst ratio is 3.3 mL g
1. The reaction conditions under the wet
dry multiphase conditions and a liquid film state are basically the same. Kariya and co-workers used the term “wet
dry multiphase condition” rather than “liquid film state” because the latter has some static nuances, as if the thin film of the liquid reactant and product always covers the entire catalyst surface. The term “wet
dry multiphase condition”, on the other hand, clearly emphasis the dynamic phase change between gas and liquid, as well as the drying process associated with the rapid desorption of the liquid reactant and product from the catalyst surface. The wet
dry multiphase condition was used in developing a spray-pulse reactor that efficiently dehydrogenate the liquid organic carrier by Pt-containing catalysts supported on thin active carbon cloth sheets and alumite plates.4,25,26 4229

dx.doi.org/10.1021/ef200829x |Energy Fuels 2011, 25, 4217–4234

Energy & Fuels According to Ichikawa,4 the spray-pulse reactor yields higher evolution rates of hydrogen that are 10
25 and 100
1600 times greater than those of conventional gas- and liquid-phase reactions, respectively. Other research groups75
78 tried to overcome the thermodynamically unfavorable dehydrogenation of liquid organic hydride by proposing the storage of hydrogen in liquid organic heterocyclic compounds. According to Clot et al.77 and Crabtree,78 the release of hydrogen is greatly favored thermodynamically in fivemembered ring species and not six, such as cyclohexane, and the dehydrogenation can be further enhanced by incorporating nitrogen atoms into the rings, as either a ring atom or a ring substituent. The incorporation of N or less satisfactory O into the hydrogen carrier facilitates the hydrogen release by decreasing the endothermicity of the dehydrogenation reaction. Bringing this concept to fruition, however, requires an extensive catalyst development effort to obtain a highly active and selective hydrogenation
dehydrogenation catalyst.78 In a recent effort, Luca et al.79 proposed organoelectrocatalysis as an alternative to heterogeneous catalysis for such a dehydrogenation reaction. They reported that the dehydrogenation of N-containing heterocyclic compounds can be performed by dehydrogenative oxidative transformation under electrochemical conditions. As a proof of principle, they have used 2,3-dichloro-5,6-dicyanobenzoquinone (DDQ) as an electrocatalyst for N-phenylbenzylamine dehydrogenation. Luca and co-workers found that DDQ can abstract two H atom equivalents from the NH
CH2 group not only stoichiometrically but also via electrochemical organocatalysis.

6. SULFUR IMPACT ON MCH DEHYDROGENATION The supported Pt/Al2O3 reforming catalyst and its bimetallic successors are usually presulfided (pre-poisoned) in commercial applications to improve their performance. Many catalytic studies tried to evaluate the effect of sulfur poisoning on the performance of Pt/Al2O3 and Pt
Re/Al2O3. In this section, the effect of sulfur on MCH dehydrogenation is addressed by illustrating its impact on three catalytic parameters: the activity, the selectivity, and the stability. The attraction of sulfur to Re is much stronger than its attraction to Pt.68 According to Van Trimpont et al.,80 the surface coverage of strongly adsorbed sulfur is twice as great on Pt
Re/ Al2O3 than on Pt/Al2O3. The weak adsorption of sulfur by Pt requires the continuous feeding of sulfur when reforming with Pt/Al2O3. The Pt
Re/Al2O3 catalyst, on the other hand, requires only an initial presulfiding to suppress hydrogenolysis. Continuous addition of sulfur to Pt
Re/Al2O3 normally results in a drastic decrease in the dehydrogenation activity and selectivity. Biloen et al.,68 Sachtler,64 and Parera and Beltramini55 reported that sulfur is selectively chemisorbed on Re atoms in the bimetallic Pt
Re catalyst to form Re
S, which divides the Pt surface and reduce the possibility of having large Pt ensembles. The impact of sulfur on the catalyst performance mainly depends upon its concentration. High concentrations of sulfur usually result in rapid catalyst deactivation. Low concentrations of sulfur, on the other hand, can effectively reduce the hydrogenolysis activity, which enhances the selectivity and stability. Because of that, many industrial practitioners keep the sulfur concentration below 1 ppm in the feed to enhance the catalytic performance without causing substantial deactivation. 6.1. Sulfur Impact on Reaction Activity. Many researchers have compared the dehydrogenation activity of Pt/Al2O3 and

REVIEW

Pt
Re/Al2O3 in the absence of sulfur. A group of researchers reported comparable activity of both catalysts, while others believed that the monometallic catalyst has higher dehydrogenation activity. For example, Bastein et al.81 stated that the catalytic activity of pure and alloyed platinum must be the same because both have similar electronic structures. Pal et al.57 also reported that Pt
Re/Al2O3 and Pt/Al2O3 have comparable activities because Re has no dehydrogenation impact (only Pt is responsible for dehydrogenation activity). Jossens and Petersen42 and Jothimurugesan et al.44 also observed very similar activities for both catalysts. Van Trimpont et al.,80 conversely, reported that the dehydrogenation activity of Pt/Al2O3 is 4 times higher than that of the Pt
Re/Al2O3 catalyst. Alhumaidan et al.52 compared the initial activity of various commercial and prototype catalysts for MCH dehydrogenation. They reported that the dehydrogenation activity of Pt/Al2O3 and Pt
Re/Al2O3 is not absolute and significantly depends upon Pt dispersion and morphology as well as reaction conditions. According to Alhumaidan et al.,52 a catalyst that is highly active at one set of operating conditions might not be that active at another set of conditions. Therefore, the activity comparison should be performed at sets of conditions of particular interest. In their study of the effect of sulfur poisoning on the activities of Pt/Al2O3 and Pt
Re/Al2O3, Jossens and Petersen42 reported that the sulfiding pretreatment significantly reduced the initial activity in both catalysts but the reduction in activity was more evident in Pt
Re/Al2O3. Pacheco and Petersen59 also reported that the activity of Pt
Re/Al2O3 is more sensitive to sulfur poising than Pt/Al2O3. These observations are consistent with the commercial reforming applications, where high-sulfur feed is normally treated with Pt/Al2O3 rather than Pt
Re/Al2O3. Van Trimpont et al.80 reported that the reduction in the dehydrogenation activity with sulfur addition is mainly attributed to the reduction in metallic accessible sites on the catalyst surface. Biloen et al.,68 however, believe that sulfiding should not significantly impact the activity of Pt
Re/Al2O3 because sulfur atoms are preferentially chemisorbed to Re atoms, which leave the Pt sites predominantly uncovered. The contradiction among the previous observations might be due to the difference in sulfur concentrations. At high concentrations of sulfur, the sulfur might cover more Pt atoms after depleting all accessible Re atoms. 6.2. Sulfur Impact on Reaction Selectivity. Catalyst selectivity in the dehydrogenation reaction is inversely proportional to the hydrogenolysis activity. Many researchers reported that the hydrogenolysis activity in Pt/Al2O3 is directly proportional to the size of Pt ensembles.54,82
84 Doolittle et al.54 indicated that an increase in Pt loading from 0.32 to 0.60 wt % substantially increased the size of metal crystallites and the yields of benzene and lighter products. Doolittle and his partners suggested that the demethylation and the hydrogenolysis reactions occur exclusively on large Pt crystallites. Coughlin et al.,67 on the other hand, associated the hydrogenolysis activity with kink sites (edge sites), which increase in number as the Pt ensemble decreases in size. Previous research efforts have also indicated that the addition of Re promotes the hydrogenolysis activity.42,44,68 According to the literature, Re has two roles in Pt catalysts: it reduces the size of Pt ensembles and continuously hydrocracks graphitic precursors. To achieve the optimum toluene selectivity in MCH dehydrogenation over Pt/ Al2O3 and Pt
Re/Al2O3, the hydrogenolysis activity of Pt and Re atoms must be partially suppressed by a sulfiding treatment. Various research studies reported higher toluene selectivity in MCH dehydrogenation over Pt
Re/Al2O3 as compared to 4230

dx.doi.org/10.1021/ef200829x |Energy Fuels 2011, 25, 4217–4234

Energy & Fuels Pt/Al2O3.42,44,54,85 Coughlin et al.66 also reported that the presence of Re in Pt reforming catalyst appreciably increased toluene selectivity; however, this enhancement in the toluene yield was not evident until after several hours of operation. According to Coughlin and co-workers, the gradual increase in toluene selectivity was associated with the continuing accumulation of carbonaceous deposits, which promote and facilitate toluene desorption (toluene migrates from Pt crystals to carbonaceous overlayers) and inhibit hydrogenolysis activity. Davis et al.53 and Doolittle et al.54 also revealed that carbonaceous overlayers work as desorption sites for the chemisorbed product and provide hydrogen exchange with reacting species. In the case of Pt/ Al2O3, the carbonaceous overlayers rapidly convert to graphitic deposits that consequently block the Pt sites and impede toluene desorption. For Pt
Re/Al2O3, on the other hand, the early accumulation of carbonaceous overlayers (first 2 h) gradually promotes toluene selectivity by partially deactivating hydrogenolysis sites. The additional accumulations of carbonaceous overlayers on Pt
Re/Al2O3 cannot easily form graphitic deposits because Re continuously hydrocracks the graphitic precursors.54,66 Many researchers reported that the selective poisoning of kink sites by sulfur appreciably suppresses hydrogenolysis activity on Pt/Al2O3.54,55,67,80 Coughlin et al.67 reported that sulfided reforming catalysts have higher initial selectivity toward toluene and lower benzene yield than the unsulfided catalysts. This initial selectivity toward toluene was attributed to the rapid and selective chemisorption of sulfur on the metal cracking sites of Pt ensembles. The unsulfided catalysts, in contrast, illustrated a slower increase in toluene selectivity because the deactivation of kink sites was achieved by carbonaceous deposits (instead of sulfur), which require time to accumulate. 6.3. Sulfur Impact on Reaction Stability. As previously indicated, the addition of Re to Pt/Al2O3 significantly enhanced the catalyst stability.46,49,54,57,60,63,64,66 Various explanations for the impact of Re on stability enhancement have been illustrated. In this section, the influence of sulfur poisoning on catalyst stability will be discussed. Many researchers reported that sulfur poisoning considerably increased catalyst stability in MCH dehydrogenation. Van Trimpont et al.,46 Doolittle et al.,54 and Chai and Kawakami49 reported that the addition of sulfur preferentially impeded hydrogenolysis reactions, which reduces the fouling rate and prolongs the catalyst life. Pal et al.57 also stated that the adsorbed sulfur in Pt
Re/Al2O3 prevents the ultimate transformation of coke precursors. Pacheco and Petersen45 further studied the effect of sulfur poisoning on the stability of Pt/Al2O3 and Pt
Re/Al2O3 and revealed a significant drop in the fouling rate, especially in the presulfided bimetallic catalyst. According to them, the specific rate of fouling (kf) in the sulfided bimetallic catalyst was 20 times less than that of the monometallic catalyst for the same pretreatment and operating conditions. They explained the drop in the fouling rate using the previously proposed multiplet fouling mechanism. They believed that the addition of sulfur deactivates the larger multiplets, which consequently reduce the initial deactivation and allow the catalyst to operate at higher activity and stability. Coughlin et al.67 investigated the influence of sulfur on MCH dehydrogenation over Pt/Al2O3 and Pt
Re/Al2O3 catalysts using thermogravimetric analysis (TGA). The TGA results indicated that the presence of sulfur significantly retarded coke deposition, especially in the bimetallic catalyst. In their study, Coughlin and co-workers also compared two sulfiding techniques: the

REVIEW

continuous sulfiding with thiophene and the presulfiding by H2S. The second technique was more effective for the bimetallic catalyst because the continuous sulfiding added an excessive amount of sulfur, which considerably retarded the dehydrogenation activity. The effect of excessive sulfur was less pronounced with the monometallic catalyst. Parera and Beltramini55 studied coke deposition on Pt/Al2O3, Pt
Re/Al2O3, and Pt
Re
S/Al2O3 catalysts using temperature-programmed oxidation (TPO). The study indicated that the coke is deposited on both metallic and acidic sites of the catalysts. At the beginning of the run, the coke deposition mainly occurred on the metallic sites. As the run proceeded, the deposition started to shift gradually toward the acidic sites. Parera and Beltramini compared the catalyst deactivation by coke deposition for the three catalysts. They found that Pt/Al2O3 was covered with coke more rapidly than Pt
Re/Al2O3 and Pt
Re
S/Al2O3. They also noticed that more coke was formed on Pt
Re
S/Al2O3 than Pt
Re/Al2O3; however, the sulfided catalyst showed more resistance to deactivation. On the basis of this deactivation behavior, Parera and Beltramini55 declared that the deactivation of the catalyst depends upon not only the amount of coke but also the distribution of coke on the catalyst surface. According to them, the addition of Re and S stabilizes the catalyst in two ways. First, the inert Re
S decreases the Pt ensemble size, which consequently reduces the hydrogenolysis activity and coke formation rate on the metallic function. Second, the presulfiding treatment reduces the deposition of coke on the metallic function by allowing coke precursors to migrate from the platinum sites to the support. Parera and Beltramini proved their argument using TPO to differentiate between coke deposited on metallic and acidic functions. The coke deposited on the metallic sites combusts at 350
370 °C, while the coke formed on acidic sites combusts at 370
550 °C. Coughlin et al.67 studied the regeneration of sulfided Pt/Al2O3 and Pt
Re/Al2O3 by cyclically exposing them to a stream of pure hydrogen. They reported that hydrogen treatment of the sulfided mono- and bimetallic catalysts has restored the dehydrogenation activity of both catalysts but could not restore the hydrogenolysis activity, which has been irreversibly poisoned by adsorbed sulfur. Guenin et al.86 reported that the unsulfided catalyst has a higher activity recovery than the sulfided catalyst. Doolittle et al.,54 on the contrary, indicated that sulfur addition has no influence on catalyst regeneration by hydrogen treatment. The previous discussion indicates that sulfur addition can have either beneficial or detrimental effects on MCH dehydrogenation. The key factor in that is the proper concentration of sulfur added to the catalyst. Low concentrations of sulfur suppress hydrogenolysis reactions and retard coke formation, while an excessive addition of sulfur might block the metallic sites and suppress the dehydrogenation activity.

7. CONCLUSION The use of liquid organic hydride in hydrogen storage has technical, economical, and environmental advantages. One of the most promising organic cycles for hydrogen storage is the MTH cycle, which consists of MCH, toluene, and hydrogen. The main obstacle that prevent the commercial establishment of this concept is related to the dehydrogenation aspect of this cycle, which is highly endothermic. Various catalysts have been developed to enhance the activity, selectivity, and stability of the dehydrogenation reaction. The results clearly suggest that Pt/Al2O3 and its 4231

dx.doi.org/10.1021/ef200829x |Energy Fuels 2011, 25, 4217–4234

Energy & Fuels Pt
Re/Al2O3 successor are the most promising catalysts. This review clearly demonstrates the considerable discrepancy in the literature with regard to the kinetics of MCH dehydrogenation over supported Pt and Pt
Re catalysts. The dispute on the poisoning mechanism and deactivation kinetics of Pt/Al2O3 and Pt
Re/Al2O3 during MCH dehydrogenation were also illustrated. The review also reveals the attempt to exceed the equilibrium limitation on MCH dehydrogenation. The beneficial and detrimental impacts of sulfur addition to MCH dehydrogenation were also covered in this review.

’ AUTHOR INFORMATION Corresponding Author

*E-mail: [email protected].

’ REFERENCES (1) Okada, Y.; Sasaki, E.; Watanabe, E.; Hyodo, S.; Nishijima, H. Development of dehydrogenation catalyst for hydrogen generation in organic chemical hydride method. Int. J. Hydrogen Energy 2006, 31 (10), 1348–1356. (2) Pradhan, A. U.; Shukla, A.; Pande, J. V.; Karmarkar, S.; Biniwale, R. B. A feasibility analysis of hydrogen delivery system using liquid organic hydrides. Int. J. Hydrogen Energy 2011, 36 (1), 680–688. (3) Giordano, N.; Cacciola, G.; Parmaliana, A. Energy storage and transmission: Platinum catalyses chemical closed-loop system. Platinum Met. Rev. 1986, 30 (4), 174–182. (4) Ichikawa, M. Organic liquid carrier for hydrogen storage. SolidState Hydrogen Storage 2008, 500–531. (5) Scherer, G. W. H.; Newson, E. Analysis of the seasonal energy storage of hydrogen in liquid organic hydrides. Int. J. Hydrogen Energy 1998, 23 (1), 19–25. (6) Scherer, G. W. H.; Newson, E.; Wokaun, A. Economic analysis of the seasonal storage of electricity with liquid organic hydrides. Int. J. Hydrogen Energy 1999, 24 (12), 1157–1169. (7) DeLuchi, M. A. Hydrogen vehicles: An evaluation of fuel storage, performance, safety, environmental impacts, and cost. Int. J. Hydrogen Energy 1989, 14 (2), 81–130. (8) Taube, M.; Rippin, D. W. T.; Cresswell, D. L.; Knecht, W. A system of hydrogen-powered vehicles with liquid organic hydrides. Int. J. Hydrogen Energy 1983, 8 (3), 213–225. (9) Taube, M.; Rippin, D.; Knecht, W.; Hakimifard, D.; Milisavljevic, B.; Gruenenfelder, N. A prototype truck powered by hydrogen from organic liquid hydrides. Int. J. Hydrogen Energy 1985, 10 (9), 595–599. (10) Grunenfelder, N. F.; Schucan, T. H. Seasonal storage of hydrogen in liquid organic hydrides: Description of the second prototype vehicle. Int. J. Hydrogen Energy 1989, 14 (8), 579–586. (11) Kariya, N.; Fukuoka, A.; Ichikawa, M. Efficient evolution of hydrogen from liquid cycloalkanes over Pt-containing catalysts supported on active carbons under “wet
dry multiphase conditions”. Appl. Catal., A 2002, 233 (1
2), 91–102. (12) Klvana, D.; Touzani, A.; Chaouki, J.; Belanger, G. Dehydrogenation of methylcyclohexane in a reactor coupled to a hydrogen engine. Int. J. Hydrogen Energy 1991, 16 (1), 55–60. (13) Newson, E.; Haueter, T.; Hottinger, P.; Von Roth, F.; Scherer, G. W. H.; Schucan, T. H. Seasonal storage of hydrogen in stationary systems with liquid organic hydrides. Int. J. Hydrogen Energy 1998, 23 (10), 905–909. (14) Biniwale, R. B.; Rayalu, S.; Devotta, S.; Ichikawa, M. Chemical hydrides: A solution to high capacity hydrogen storage and supply. Int. J. Hydrogen Energy 2008, 33 (1), 360–365. (15) Rahimpour, M. R.; Vakili, R.; Pourazadi, E.; Bahmanpour, A. M.; Iranshahi, D. Enhancement of hydrogen production via coupling of MCH dehydrogenation reaction and methanol synthesis process by using thermally coupled heat exchanger reactor. Int. J. Hydrogen Energy 2011, 36 (5), 3371–3383.

REVIEW

(16) Shukla, A. A.; Gosavi, P. V.; Pande, J. V.; Kumar, V. P.; Chary, K. V. R.; Biniwale, R. B. Efficient hydrogen supply through catalytic dehydrogenation of methylcyclohexane over Pt/metal oxide catalysts. Int. J. Hydrogen Energy 2010, 35 (9), 4020–4026. (17) Shukla, A.; Karmakar, S.; Biniwale, R. B. Hydrogen delivery through liquid organic hydrides: Considerations for a potential technology. Int. J. Hydrogen Energy 2011, DOI: 10.1016/j.ijhydene.2011.04.107. (18) Gretz, J.; Baselt, J. P.; Ullmann, O.; Wendt, H. The 100 MW euro-Quebec hydro-hydrogen pilot project. Int. J. Hydrogen Energy 1990, 15 (6), 419–424. (19) Gretz, J.; Drolet, B.; Kluyskens, D.; Sandmann, F.; Ullmann, O. Status of the hydro-hydrogen pilot project (EQHHPP). Int. J. Hydrogen Energy 1994, 19 (2), 169–174. (20) Sultan, O.; Shaw, H. Study of Automotive Storage of Hydrogen Using Recyclable Liquid Chemical Carriers. Technical Report TEC-75/ 003, 1975; http://www.osti.gov/energycitations/product.biblio.jsp? osti_id=5000657. (21) Zhang, C.; Liang, X.; Liu, S. Hydrogen production by catalytic dehydrogenation of methylcyclohexane over Pt catalysts supported on pyrolytic waste tire char. Int. J. Hydrogen Energy 2011, 36, 8902–8907. (22) Wang, Y.; Shah, N.; Huffman, G. Pure hydrogen production by partial dehydrogenation of cyclohexane and methylcyclohexane over nanotube-supported Pt and Pd catalysts. Energy Fuels 2004, 18 (5), 1429–1433. (23) Cacciola, G.; Giordano, N.; Restuccia, G. Cyclohexane as a liquid phase carrier in hydrogen storage and transport. Int. J. Hydrogen Energy 1984, 9 (5), 411–419. (24) Touzani, A.; Klvana, D.; Belanger, G. Dehydrogenation reactor for a vehicle equipped with a hydrogen engine: A simulation study. Int. J. Hydrogen Energy 1984, 9 (11), 929–936. (25) Kariya, N.; Fukuoka, A.; Utagawa, T.; Sakuramoto, M.; Goto, Y.; Ichikawa, M. Efficient hydrogen production using cyclohexane and decalin by pulse-spray mode reactor with Pt catalysts. Appl. Catal., A 2003, 247 (2), 247–259. (26) Biniwale, R. B.; Ichikawa, M. Thermal imaging of catalyst surface during catalytic dehydrogenation of cyclohexane under spraypulsed conditions. Chem. Eng. Sci. 2007, 62 (24), 7370–7377. (27) Newson, E.; Truong, T. B.; Hottinger, P.; Roth, F. V.; Schucan, T. H. H. Optimisation of seasonal energy storage in stationary systems with liquid hydrogen carriers, decalin and methylcyclohexane. Proceedings of the 12th World Hydrogen Energy Conference; Buenos Aires, Argentina, June 21
26, 1998; Vol. 2, pp 935
942. (28) Rahimpour, M. R.; Mirvakili, A.; Paymooni, K. Hydrogen as an energy carrier: A comparative study between decalin and cyclohexane in thermally coupled membrane reactor in gas-to-liquid technology. Int. J. Hydrogen Energy 2011, 36, 6970–6984. (29) Shinohara, C.; Kawakami, S.; Moriga, T.; Hayashi, H.; Hodoshima, S.; Saito, Y.; Sugiyama, S. Local structure around platinum in Pt/C catalysts employed for liquid-phase dehydrogenation of decalin in the liquid-film state under reactive distillation conditions. Appl. Catal., A 2004, 266 (2), 251–255. (30) Hodoshima, S.; Arai, H.; Saito, Y. Liquid-film-type catalytic decalin dehydrogeno-aromatization for long-term storage and longdistance transportation of hydrogen. Int. J. Hydrogen Energy 2003, 28 (2), 197–204. (31) Hodoshima, S.; Arai, H.; Takaiwa, S.; Saito, Y. Catalytic decalin dehydrogenation/naphthalene hydrogenation pair as a hydrogen source for fuel-cell vehicle. Int. J. Hydrogen Energy 2003, 28 (11), 1255–1262. (32) Hodoshima, S.; Takaiwa, S.; Shono, A.; Satoh, K.; Saito, Y. Hydrogen storage by decalin/naphthalene pair and hydrogen supply to fuel cells by use of superheated liquid-film-type catalysis. Appl. Catal., A 2005, 283 (1
2), 235–242. (33) Sebastian, D.; Bordeje, E. G.; Calvillo, L.; Lazaro, M. J.; Moliner, R. Hydrogen storage by decalin dehydrogenation/naphthalene hydrogenation pair over platinum catalysts supported on activated carbon. Int. J. Hydrogen Energy 2008, 33 (4), 1329–1334. (34) Richardson, J. Principles of Catalyst Development; Plenum Press: New York, 1989; p 288. 4232

dx.doi.org/10.1021/ef200829x |Energy Fuels 2011, 25, 4217–4234

Energy & Fuels (35) Corma, A.; Cid, R.; Agudo, A. Catalyst decay in the kinetics of methylcyclohexane dehydrogenation over Pt
NaY zeolite. Can. J. Chem. Eng. 1979, 57, 638. (36) García de la Banda, J. F.; Corma, A.; Melo, F. V. Dehydrogenation of methylcyclohexene on a PtNaY catalyst. Study of kinetics and deactivation. Appl. Catal. 1986, 26, 103–121. (37) Sinfelt, J. H. The turnover frequency of methylcyclohexane dehydrogenation to toluene on a Pt reforming catalyst. J. Mol. Catal. A: Chem. 2000, 163 (1
2), 123–128. (38) Sinfelt, J.; Hurwitz, H.; Shulman, R. Kinetics of methylcyclohexane dehydrogenation over Pt
Al2O3. J. Phys. Chem. 1960, 64 (10), 1559–1562. (39) Hawthorn, R. D.; Ackerman, G. H.; Nixon, A. C. A mathmatical model of a packed-bed heat-exchanger reactor for dehydrogenation of methylcyclohexane: Comparison of predictions with experimental results. AIChE J. 1968, 14 (1), 69–76. (40) Wolf, E. E.; Petersen, E. E. Kinetics of deactivation of a reforming catalyst during methylcyclohexane dehydrogenation in a diffusion reactor. J. Catal. 1977, 46 (2), 190–203. (41) Jossens, L. W.; Petersen, E. E. Fouling of a platinum reforming catalyst accompanying the dehydrogenation of methyl cyclohexane. J. Catal. 1982, 73 (2), 377–386. (42) Jossens, L. W.; Petersen, E. E. Fouling of a platinum
rhenium reforming catalyst using model reforming reactions. J. Catal. 1982, 76 (2), 265–273. (43) Andreev, A.; Shopov, D.; Kiperman, S. Kinetics of cyclohexane dehydrogenation in gradientless system. Kinet. Catal. 1966, 7, 941. (44) Jothimurugesan, K.; Bhatla, S.; Srivastava, R. D. Kinetics of dehydrogenation of methylcyclohexane over platinum
rhenium
alumina catalyst in the presence of added hydrogen. Ind. Eng. Chem. Fundam. 1985, 24, 433–438. (45) Pacheco, M. A.; Petersen, E. E. Reaction kinetics of methylcyclohexane dehydrogenation over a sulfided Pt + Re/Al2O3 reforming catalyst. J. Catal. 1985, 96 (2), 507–516. (46) Van Trimpont, P. A.; Marin, G. B.; Froment, G. F. Kinetics of methylcyclohexane dehydrogenation on sulfided commercial platinum/ alumina and platinum
rhenium/alumina catalyst. Ind. Eng. Chem. Fundam. 1986, 25, 544–553. (47) Van Trimpont, P. A.; Marin, G. B.; Froment, G. F. Kinetics of the reforming of C7 hydrocarbons on a commercial PtRe/Al2O3 catalyst. Appl. Catal. 1986, 24 (1
2), 53–68. (48) El-Sawi, M.; Infortuna, F. A.; Lignola, P. G.; Parmaliana, A.; Frusteri, F.; Giordano, N. Parameter estimation in the kinetic model of methylcyclohexane dehydrogenation on a Pt
Al2O3 catalyst by sequential experiment design. Chem. Eng. J. 1989, 42 (3), 137–144. (49) Chai, M.; Kawakami, K. Kinetic model and simulation for catalyst deactivation during dehydrogenation of methylcyclohexane over commercial Pt, PtRe and presulfided PtRe
Al2O3 catalyst. J. Chem. Technol. Biotechnol. 1990, 51, 335–345. (50) Tsakiris, D. Catalytic production of hydrogen from liquid organic hydride. Ph.D. Thesis, The University of Manchester, Manchester, U.K., 2008. (51) Alhumaidan, F. Hydrogen storage in liquid organic hydride: Producing hydrogen catalytically from methylcyclohexane. Ph.D. Thesis, The University of Manchester, Manchester, U.K., 2008. (52) Alhumaidan, F.; Cresswell, D.; Garforth, A. Kinetic model of the dehydrogenation of methylcyclohexane over monometallic and bimetallic Pt catalysts. Ind. Eng. Chem. Res. 2011, 50, 2509–2522. (53) Davis, S. M.; Zaera, F.; Somorjai, G. A. The reactivity and composition of strongly adsorbed carbonaceous deposits on platinum. Model of the working hydrocarbon conversion catalyst. J. Catal. 1982, 77 (2), 439–459. (54) Doolittle, W. J.; Skoularikis, N. D.; Coughlin, R. W. Reactions of methylcyclohexane and n-heptane over supported Pt and Pt
Re catalysts. J. Catal. 1987, 107 (2), 490–502. (55) Parera, J. M.; Beltramini, J. N. Stability of bimetallic reforming catalysts. J. Catal. 1988, 112 (2), 357–365. (56) Ali, L. I.; Ali, A.-G. A.; Aboul-Fotouh, S. M.; Aboul-Gheit, A. K. Dehydrogenation of cyclohexane on catalysts containing noble metals

REVIEW

and their combinations with platinum on alumina support. Appl. Catal., A 1999, 177 (1), 99–110. (57) Pal, A. K.; Bhowmick, M.; Srivastava, R. D. Deactivation kinetics of platinum
rhenium reforming catalyst accompanying the dehydrogenation of methylcyclohexane. Ind. Eng. Chem. Process Des. Dev. 1986, 25, 236–241. (58) Pacheco, M. A.; Petersen, E. E. On a general correlation for catalyst fouling. J. Catal. 1984, 86 (1), 75–83. (59) Pacheco, M. A.; Petersen, E. E. On the development of a catalyst fouling model. J. Catal. 1984, 88 (2), 400–408. (60) Alhumaidan, F.; Cresswell, D.; Garforth, A. Long-term deactivation of supported Pt catalysts in the dehydrogenation of methylcyclohexane to toluene. Ind. Eng. Chem. Res. 2010, 49, 9764–9770. (61) Engels, S.; Lehmann, L. Investigation on metal catalysts. XIV. Physico-chemcial and catalytic properties of palladium
rehium catalysts. Z. Anorg. Allg. Chem. 1977, 431 (1), 201–206. (62) Leprince, P. Conversion Process; Institute Francais du Petrole: Paris, France, 2001. (63) Bertolacini, R. J.; Pellet, R. J. The function of rhenium in bimetallic reforming catalysis. Stud. Surf. Sci. Catal. 1980, 6, 73–77. (64) Sachtler, W. M. H. Selectivity and rate of activity decline of bimetallic catalysts. J. Mol. Catal. 1984, 25 (1
3), 1–12. (65) Somorjai, G. A.; Blakely, D. W. Mechanism of catalysis of hydrocarbon reactions by platinum surfaces. Nature 1975, 258, 580–583. (66) Coughlin, R. W.; Kawakami, K.; Hasan, A. Activity, yield patterns, and coking behavior of Pt and PtRe catalysts during dehydrogenation of methylcyclohexane: I. In the absence of sulfur. J. Catal. 1984, 88 (1), 150–162. (67) Coughlin, R. W.; Hasan, A.; Kawakami, K. Activity, yield patterns, and coking behavior of Pt and PtRe catalysts during dehydrogenation of methylcyclohexane: II. Influence of sulfur. J. Catal. 1984, 88 (1), 163–176. (68) Biloen, P.; Helle, J. N.; Verbeek, H.; Dautzenberg, F. M.; Sachtler, W. M. H. The role of rhenium and sulfur in platinum-based hydrocarbon-conversion catalysts. J. Catal. 1980, 63 (1), 112–118. (69) Ali, J. K.; Newson, E. J.; Rippin, D. W. T. Exceeding equilibrium conversion with a catalytic membrane reactor for the dehydrogenation of methylcyclohexane. Chem. Eng. Sci. 1994, 49 (13), 2129–2134. (70) Ali, J. K.; Baiker, A. Measuring conversion in dehydrogenation reactions with the aid of a Pd
Ag membrane. Ind. Eng. Chem. Res. 1995, 34, 3166–3169. (71) Ali, J. K.; Baiker, A. Dehydrogenation of methylcyclohexane to toluene in a pilot-scale membrane reactor. Appl. Catal., A 1997, 155 (1), 27–39. (72) Schildhauer, T.; Newson, E.; Muller, S. The equilibrium constant for the methylcyclohexane
toluene system. J. Catal. 2001, 198 (2), 355–358. (73) Ferreira-Aparicio, P.; Rodriguez-Ramos, I.; Guerrero-Ruiz, A. Pure hydrogen production from methylcyclohexane using a new high performance membrane reactor. Chem. Commun. 2002, 2082–2083. (74) Ferreira-Aparicio, P.; Rodriguez-Ramos, I.; Guerrero-Ruiz, A. On the performance of porous vycor membranes for conversion enhancement in the dehydrogenation of methylcyclohexane to toluene. J. Catal. 2002, 212 (2), 182–192. (75) Pez, G. P.; Scott, A. R.; Cooper, A. C.; Cheng, H. Hydrogen storage by reversible hydrogenation of π-conjugated substrates. U.S. Patent 7,101,530, 2006. (76) Pez, G. P.; Scott, A. R.; Cooper, A. C.; Cheng, H.; Wilhelm, F. C.; Abdourazak, A. H. Hydrogen storage by reversible hydrogenation of π-conjugated substrates. U.S. Patent 7,351,395, 2008. (77) Clot, E.; Eisenstein, O.; Crabtree, R. Computational structure
activity relationships in H2 storage—How placement of N atoms affects release temperature in organic liquid storage materials. Chem. Commun. 2007, 2231–2233. (78) Crabtree, R. Hydrogen storage in liquid organic heterocycles. Energy Environ. Sci. 2008, 1, 134–138. (79) Luca, O. R.; Wang, T.; Konezny, S. J.; Batista, V. S.; Crabtree, R. H. DDQ as an electrocatalyst for amine dehydrogenation, a model 4233

dx.doi.org/10.1021/ef200829x |Energy Fuels 2011, 25, 4217–4234

Energy & Fuels

REVIEW

system for virtual hydrogen storage. New J. Chem. 2011, 35 (5), 998–999. (80) Van Trimpont, P. A.; Marin, G. B.; Froment, G. F. Activities and selectivities for reforming reactions on unsulfided and sulfided commercial platinum and platinum
rhenium catalysts. Appl. Catal. 1985, 17 (1), 155–160. (81) Bastein, A. G. T. M.; Toolenaar, F. J. C. M.; Ponec, V. Adsorption of carbon monoxide on platinum alloys: An infrared investigation. J. Catal. 1984, 90 (1), 88–95. (82) Sinfelt, J. H.; Carter, J. L.; Yates, D. J. C. Catalytic hydrogenolysis and dehydrogenation over copper
nickel alloys. J. Catal. 1972, 24 (2), 283–296. (83) Ponec, V.; Sachtler, W. M. H. The reactions between cyclopentane and deuterium on nickel and nickel
copper alloys. J. Catal. 1972, 24 (2), 250–261. (84) Dautzenberg, F.; Helle, J.; Biloen, P.; Sachtler, W. Conversion of n-hexane over monofunctional supported and unsupported PtSn catalysts. J. Catal. 1980, 63 (1), 119–128. (85) Cid, R.; Reyes, P.; Pecchi, G. Catalytic activity for dehydrogenation of methylcyclohexane using a platinum
rhenium catalyst supported on zeolite NaY. J. Chem. Res. 1986, 1, 64–65. (86) Guenin, M.; Breysse, M.; Frety, R.; Tifouti, K.; Marecot, P.; Barbier, J. Resistance to sulfur poisoning of metal catalysts: Dehydrogenation of cyclohexane on Pt/Al2O3 catalysts. J. Catal. 1987, 105 (1), 144–154.

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dx.doi.org/10.1021/ef200829x |Energy Fuels 2011, 25, 4217–4234