Conversion of Methane and Carbon Dioxide to Higher Value Products

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Conversion of Methane and Carbon Dioxide to Higher Value Products Vesna Havran, Milorad P. Dudukovic, and Cynthia S. Lo* Department of Energy, Environmental and Chemical Engineering, Washington University in Saint Louis, One Brookings Drive, St. Louis, Missouri 63130, United States ABSTRACT: In this manuscript we address the efforts reported in the literature on direct conversion of methane and carbon dioxide into higher value products. The abundance of these two greenhouse gases makes them potentially useful for chemical syntheses. However, strong intramolecular bonds render both molecules chemically inert and thermodynamically stable. Thus, significant energy inputs as well as properly designed catalytic systems are required for their reactions. In addition to traditional catalytic processes, new alternative routes—such as photocatalytic conversion and dielectric barrier discharges—have been suggested. Figures of merit for comparison of various approaches have not yet been reported and are overdue.

1. INTRODUCTION Methane and carbon dioxide are well-known greenhouse gases. Although the amounts of methane in the atmosphere are lower than those of carbon dioxide, methane’s global warming potential (GWP) is approximately 25 times higher than that of carbon dioxide.1 Large amounts of methane are widely available in nature in the form of natural gas while substantial amounts of carbon dioxide are generated by human activity. The abundance of these two gases makes them attractive raw materials for fuels and chemical synthesis. Currently, proven world natural gas reserves are estimated to 6609 trillion cubic feet or around 187 trillion cubic meters according to the latest reports (U.S. Energy Information Administration, Figure 1).2 Large amounts of natural gas are mainly found in remote areas (Table 1), thus hindering its full exploitation due to cost ineffective gas transportation. This challenge increases the need for onsite gas-to-liquids conversion.37 Natural gas liquefied by refrigeration can be transported by pipelines or shipped by tankers. However, to distribute the gas by these pipelines, compression to ca. 80 atm is needed3 and sometimes pipelines may not be available for distant markets.5 The conversion of these two gases at mild conditions for the production of highly valuable chemicals and/or clean fuels is a subject of great importance in C1 chemistry. Developing an effective process for such conversion would help address environmental issues, by lowering the levels of CO2, and would reduce the need for indirect routes of fuel production from methane via syngas. It has been suggested that a direct reaction of CO2 and CH4 would be an ideal redox process.8,9 Despite great scientific attention to this issue, no breakthrough technology has emerged so far as the process of direct conversion suffers from both thermodynamic and kinetic limitations.7 Due to the variety of possible chemical routes and broadness of the topic, the focus here will be only on direct reactions between CH4 and CO2. We briefly review the key issues below. 2. LOW REACTIVITY AND THERMODYNAMIC LIMITATIONS As already mentioned, chemical utilization is one of the possible effective ways to decrease current CO2 emissions. Despite the large number of known processes for synthesis of r 2011 American Chemical Society

chemicals from carbon dioxide, only a few are actually implemented on a large scale.10 In chemical manufacturing, the current largest use of CO2 is in the synthesis of urea (a widely used fertilizer), in the production of salicylic acid (which is found in pharmaceuticals), and in cyclic organic carbonates (Table 2). Carbon dioxide has also been used for refrigeration, beverage carbonation, and dry cleaning, and in air conditioners, fireextinguishers, separation techniques, and water treatment, etc. As reported by the International Energy Agency (IEA) the total anthropogenic CO2 emissions in 2007 were around 29 109 metric tons.11 Electricity and heat generation combined with the use of hydrocarbon-based fuels in transportation have been the largest contributors of CO2, creating almost two-thirds of global emissions, as shown in Figure 2. According to IEA statistics,11 coal has been the dominant source for world electricity and heat generation at 41%, while the share of gas was 23%. The total amount of carbon dioxide used in industry is approximately 115 106 metric tons per year.12,13 Major industrial processes that utilize CO2 as a raw material are listed in Table 2. Its utilization as a technological fluid, where carbon dioxide is not chemically converted and thus can be recovered at the end, is estimated at 18 106 metric tons per year.13 The generation of energy used for CO2 transformation, if based on hydrocarbon raw materials, ironically still produces large amounts of CO2. In addition the resulting organic chemicals in which CO2 is incorporated as the result of these transformations release CO2 at the end of their use.14 However, despite the fact that, currently, the usage of CO2 in the chemical industry cannot reduce significantly the global CO2 levels, it is believed that the full potential of the fixation of CO2 into value added products has not yet been completely explored. The direct use of nonhydrocarbon based energy sources for reduction of CO2 would change the above perspective and allow recycling of carbon dioxide via chemicals and liquid fuels. Thus, the further

Received: January 4, 2011 Accepted: May 9, 2011 Revised: May 6, 2011 Published: May 23, 2011 7089

dx.doi.org/10.1021/ie2000192 | Ind. Eng. Chem. Res. 2011, 50, 7089–7100

Industrial & Engineering Chemistry Research

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Figure 1. World proven reserves of natural gas2 ( 1012 m3).

Table 1. Location of Natural Gas Reserves6 ( 1012 m3) type of location easy onshore zones

1960

1970

1990 60

15.8

27.5

offshore

1.6

4.5

25

Arctic and Siberia

0.1

7.5

42

0.5

2

world

17.5

40.0

129

share of difficult zones (%)

10

31

other difficult onshore

53

Table 2. Main Uses of Carbon Dioxide in the Chemical Industry13 major industrial processes

approximate amount of CO2

that utilize CO2 as

used in the process per

raw material

year (metric tons)

urea

70 106

inorganic carbonates and pigments methanol

30 106 6 106

salicylic acid

20 103

development of industrial processes that are utilizing CO2 for high-demand products is of importance.12,14 Of the processes listed in Table 2, the syntheses of urea and salicylic acid are thermal reactions while the other processes are catalytic. The production of urea is a two-step process where ammonia and carbon dioxide react to form ammonium carbamate which is then dehydrated to urea. The urea synthesis reactor operates at around 453483 K and at 150 atm pressure. Salicylic acid has been produced by treating sodium phenolate with carbon dioxide at a temperature of 390 K and a high pressure of about 100 atm; this is also known as the KolbeSchmitt reaction. In summary, to date, only a few compounds have been successfully synthesized from CO2 using heterogeneous catalytic routes, since most industrial processes utilize the more reactive but highly toxic carbon monoxide (CO) as a building block to produce higher chain hydrocarbons and liquid fuels. Processes for potential future utilization of CO2 that have not yet been

Figure 2. World CO2 emissions by sector for 2007.11

scaled up for commercialization are dependent on the development of new catalytic systems.13 Environmental concerns, as well as the recurring energy availability issues, raise enormous scientific interest in methane and carbon dioxide activation. Considerable literature has been devoted to the possibilities of CO2 conversion10,13,1519 as well as to the catalysis for CO2 conversion.2022 According to their energy requirements, reactions of CO2 can be divided into two groups:10 • Reactions that use the entire molecule, such as fixation onto an organic substrate in which that substrate donates the needed energy. These reactions do not need much additional energy and hence occur at lower temperatures (240400 K). • Reactions that convert CO2 into another C1 molecule or Cn molecules. These reduction reactions use hydrogen, electrons, or heat as the energy source. Because they require a large amount of energy, they occur at high temperatures (6001000 K). Carbon dioxide possesses significant thermodynamic stability. The molecule is linear with a bond strength of 532 kJ/ mol. Its Gibbs free energy of formation has a large negative value (ΔGf° = 394.6 kJ/mol), which contributes to the high inertness of CO2 and renders its reactions energetically unfavorable (Table 3). 7090

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Table 3. Enthalpy and Gibbs Free Energy Changes of Various CO2 Reactions23 ΔHr° (kJ/mol)

ΔGr° (kJ/mol)

exothermic reactions CO2(g) þ H2(g) f HCOOH(l)

31.0

þ34.3

CO2(g) þ 2H2(g) f HCHO(g) þ H2O(l)

11.7

þ46.6

CO2(g) þ 3H2(g) f CH3OH(l) þ H2O(l)

137.8

10.7

39.3

þ85.3

2CO2(g) þ H2(g) f (COOH)2(l) 2CO2(g) þ 6H2(g) f CH3OCH3(g) þ 3H2O(l)

264.9

38.0

CO2(g) þ H2(g) þ CH3OH(l) f HCOOCH3(l) þ H2O(l)

31.8

þ25.8

CO2(g) þ H2(g) þ CH3OH(l) f CH3COOH (l) þ H2O(l)

135.4

63.6

CO2(g) þ 3H2(g) þ CH3OH(l) f C2H5OH(l) þ 2H2O(l) CO2(g) þ H2(g) þ NH3(g) f HCONH2(l) þ H2O(l)

221.6 103.0

88.9 þ7.2

CO2(g) þ CH4(g) f CH3COOH(l)

13.3

þ58.1

CO2(g) þ CH4(g) þ H2(g) f CH3CHO(l) þ H2O(l)

14.6

þ74.4

CO2(g) þ CH4(g) þ 2H2(g) f (CH3)2CO(l) þ H2O(l)

70.5

þ51.2

223.6

115.0

CO2(g) þ C2H2(g) þ H2(g) f CH2dCHCOOH(l) CO2(g) þ C2H4(g) f CH2dCHCOOH(l)

49.1

þ26.2

CO2(g) þ C2H4(g) þ H2(g) f C2H5COOH(l)

166.6

56.6

CO2(g) þ C2H4(g) þ 2H2(g) f C2H5CHO(l) þ H2O(l) CO2(g) þ C6H6(l) f C6H5COOH(l)

171.1 21.6

44.4 þ30.5

6.6

þ46.9

CO2(g) þ CH2dCH2(g) f CH2CH2O(l) þ CO(g)

þ152.9

þ177.3

CO2(g) þ C(s) f 2CO(g)

þ172.6

þ119.9

3CO2(g) þ CH4(g) f 4CO (g) þ 2H2O(l)

þ235.1

þ209.2

CO2(g) þ C6H5OH(l) f m-C6H4(OH)COOH(l) endothermic reactions

CO2(g) þ CH4(g) f 2CO(g) þ 2H2(g)

þ247.5

þ170.8

CO2(g) þ 2CH4(g) f C2H6(g) þ CO(g) þ H2O(l) 2CO2(g) þ 2CH4(g) f C2H4(g) þ 2CO(g) þ 2H2O(l)

þ58.8 þ189.7

þ88.0 þ208.3

CO2(g) þ C2H4(g) f C2H4O(g) þ CO(g)

þ178.0

þ176.0

Besides requiring high energy inputs, the reduction reactions of CO2 also need an effectively designed catalytic system that lowers the kinetic activation energy barriers. A review of the current most promising CO2 reduction processes, as well as interactions of CO2 with metal surfaces and its adsorption on metal oxides, has been presented in Aresta.10 Although much is known about the mechanisms of CO reduction, elucidation of reaction pathways of the CO2 molecule on heterogeneous catalysts still remains a great challenge. Furthermore, opportunities are expanding for the use of supercritical (sc) CO2 (the state at 304 K and 72.8 atm) as an efficient and environmentally friendly reaction medium. The replacement of conventional and potentially hazardous solvents by sc CO2 offers numerous advantages: avoids the flammability, toxicity, and disposal costs, as well as the inefficient recovery and recycling related to such solvents. Detailed reviews of current and emerging applications, as well as the limitations to industrial practices, are given elsewhere in the literature.2428 Applications are broad: from the synthesis of pharmaceuticals, catalytic hydrogenation and hydroformylation reactions, carboxylation, polymerizations, etc. Supercritical CO2 has also been shown to be an efficient reaction medium for the CH bond activation in rhodium-catalyzed methane carbonylation.29 Jessop et al.30 reported that the Ru (II)phosphine complex catalyst enhanced hydrogenation of sc CO2. Although several researchers24,30,31 recognized this very attractive idea of using sc CO2 as a solvent and at the same time as the reactantC1 building block, further

Table 4. Bond Energies in Various Hydrocarbons at 298 K1 CH bond

E (kJ/mol)

CC bond

E (kJ/mol)

CH3H

434 ( 6

CH3CH3

368 ( 7

C2H5H

412 ( 6

CH3C2H5

357 ( 8

(CH3)3CH

387 ( 8

CH3C(CH3)3

C6H5H

460 ( 10

CH3C6H5

344 ( 9 417 ( 11

progress has been hindered by the lack of suitable catalysts and incomplete understanding of possible reaction pathways.24 The methane molecule, like carbon dioxide, is also very stable and symmetrical, without any functional groups, magnetic moment, or polar distribution that would assist chemical attacks and enhance its reactivity.7 The molecule has a tetrahedral geometry with four equivalent CH bonds that are very strong (434 kJ/ mol). The activation of this CH bond in the gas phase requires high temperatures at which free radical reactions are dominating.7 Furthermore, it is difficult to avoid complete oxidation32 because the energy required for breaking the CH bond in possible products is often lower (Table 4) than that in methane,1 making the target products more reactive. The main pathways for methane conversion have been classified into three main groups:9,33 • indirect routes to get syngas (steam reforming, dry reforming, and partial oxidation) • direct coupling (oxidative coupling, two-step polymerization) 7091



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Table 5. Change of Gibbs Free Energy for Various Reactions of Methane1 reaction

chemical equation

ΔGr° (kJ/mol)

(1) methane only pyrolysis

CH4 f C þ2H2

nonoxidative coupling of methane

2CH4 f C2H6 þH2

þ50.7

aromatization

6CH4 f C6H6 þ9H2

þ434

total oxidation

CH4 þ 2O2 f CO2 þ2H2O

801

oxidative coupling of methane

4CH4 þ O2 f 2C2H6þ2H2O

320

2CH4 þ O2 f C2H4þ2H2O

288

partial oxidation of methane partial oxidation of methane

2CH4 þ O2 f 2CH3OH 2CH4 þ O2 f 2CO þ4H2

223 173

partial oxidation of methane

2CH4þO2 f 2HCHO þ 2H2

104

methane to methanol

CH4 þ H2O f CH3OH þH2

þ117

steam reforming of methane

CH4 þ H2O f COþ3H2

þ142

watergas shift reaction

CO þ H2O f CO2 þ H2

steam reforming of methane þ watergas shift reaction

CH4 þ 2H2O f CO2 þ 4H2

þ114

methane to acetic acid methane to acetone

CH4 þ CO2 f CH3COOH 2CH4 þ CO2 f CH3COCH3 þ H2O

þ71.1 þ115

dry reforming of methane

CH4 þ CO2 f 2CO þ 2H2

þ171

2CH4þNH3þ2H2OfH2NCH2COOHþ5H2

þ204

þ68.6

(2) methane and oxygen

(3) methane and water

28.6

(4) methane and carbon dioxide

(5) methane and aqueous ammonia methane to amino acids

• direct conversion to oxygenates (methanol, aldehyde etc.), chlorides, or HCN As evident from the values of Gibbs free energies listed in Table 5, except for methane oxidations, most of the reactions involving methane are thermodynamically unfavorable. Besides the large amount of energy that is necessary, a suitable catalyst that would decrease the activation energy1 and help overcome reaction potential energy barrier limits has yet to be developed. Despite the considerable research effort in this area, no readily applicable solution has emerged that results in satisfactory yields and/or selectivities. Sintering of the catalyst and formation of coke or cracking products due to extensive CH and CC bond rupture, as well as increased nonselective combustion of alkanes, are all known difficulties in high-temperature alkane conversions. Developing an effective catalytic process that may be scaled up for industrial purposes remains a great challenge for catalysis.

3. METHANE ACTIVATION Successful catalyst design for reactions of methane with carbon dioxide cannot be realized without understanding of the activation of methane. It has been suggested that transition metal centers play a particularly important role in the controlled activation of the CH bond of methane in both homogeneous and heterogeneous catalysis.34 In heterogeneous reactions, studies of interaction of methane with metals have shown that adsorption of methane has been characterized as dissociative chemisorptions resulting in H2 evolution and adsorbed CHx species (where x = 0, 1, 2, or 3 depending on the nature of the metal). It has been reported that above 400 K, dehydrogenation of CHx species becomes fast and that elemental carbon is more

Figure 3. Isothermal two-step homologation process.35

stable forming stronger bonds with the metal surface.35 Studies have shown that, depending on the reaction temperature, three forms of deposited carbon can exist on the metal surface: CR reactive, mobile, or chemisorbed carbon; Cβ surface amorphous carbon; and Cγ inactive, strongly bonded graphitic carbon.35 Two research groups, Amariglio’s in France and van Santen’s in The Netherlands, independently made a considerable contribution to the understanding of the mechanisms of direct methane conversion under nonoxidative conditions and at moderate temperature.33,3539 Although their work mainly involved the study of two-step homologation of methane, their observations and results on the reactivity of adsorbed CHx species are of great importance for a wider group of reactions. The idea of Amariglio et al.38 was to overcome the thermodynamic limit and enable isothermal two step conversion by continuous removal of the hydrogen from the gas phase. This creates a full coverage of the metal surface with different CHx species leading ultimately to their recombination and homologation (Figure 3). The clear advantage 7092


Industrial & Engineering Chemistry Research of moderate temperature in the first step is very little formation of Cγ species that causes catalyst poisoning. Van Santen and co-workers theoretically investigated the reaction path for the recombination of surface CHx species on Rh and Pd clusters of 40 atoms. They concluded that the strength of the metalcarbon bond and the surface morphology are the controlling effects that determine the behavior and reactivity of CHx surface species.33 The strong metalcarbon bond leads to carbide formation, and the weak one results in hydrocarbon species. Various factors such as structure of the metal, support, metalsupport interactions, and temperature and pressure all affect the CHx species. Rasko and Solymosi40,41 examined the reactivity of adsorbed CH3 species, produced by the high-temperature pyrolysis of azomethane, toward CO2 on a Rh/SiO2 and TiO2 catalyst. Their results strongly support the idea that during the dry reforming of methane over supported Rh, the CHx fragments formed in the decomposition of methane do not decompose to carbon, but rather react with CO2. However, no activation of methane on TiO2 at 300 K has been achieved even with different pretreatments of the oxide catalyst. It has been suggested that CHx surface fragments tend to preferentially occupy adsorption sites that complete carbon tetravalency.42,43 Bradford and Vannice43 described the multistep decomposition of methane on metal surfaces in the following way: CH4 þ 2M T CH3 —M þ H—M CH3 —M þ 2M T CH2 —M2 þ H—M CH2 —M2 þ 2M T CH—M3 þ H—M CH—M3 þ 2M T C—M4 þ H—M where Mn represents an ensemble of n surface metal atoms. That renders the formation of CHx species structure sensitive. It has already been known that CH bond activation in alkanes, on well-defined surfaces, is structure sensitive as well as that coordinatively unsaturated surface atoms lead to higher methane sticking and dissociation rates than atoms on close-packed surfaces.44 Wei and Iglesia4446 studied thoroughly the mechanism and site requirements for activation and chemical conversion of methane on noble metal: Pt, Ir, Rh, Ru clusters supported on various metal oxides (ZrO2, γ-Al2O3, ZrO2CeO2). The results of their investigation of the kinetics of three different reactions of methane (methane decomposition, methane reforming with H2O, and so-called “dry reforming”—with CO2) showed no significant difference among turnover rates, rate constants, and activation energies for these three reactions. This kinetic insensitivity to coreactants and the first order CH4 kinetic rate dependence lead to the main conclusion that CH bond activation is the only kinetically relevant step. This implies that reactions of CO2 or H2O with methane derived chemisorbed carbon and CHx species are fast and kinetically irrelevant. Furthermore, since methane turnover rates increased with increasing Pt dispersion, Wei and Iglesia44 confirmed that coordinatively unsaturated Pt surface atoms, prevailing in small crystallites, are indeed more active than atoms on the low-index surfaces prevalent in larger crystallites. Those surface atoms with fewer Pt neighbors, by binding CHx and H species more strongly

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and by stabilizing the activated complexes involved in the formation of these intermediates, lower CH bond activation energies.44 Results also showed that methane reaction rates depend only on the size of Pt clusters and not on the identity of the support implying the irrelevance of the effect of the support on the overall methane reaction rates. Since their work showed that the only kinetically relevant step in all three reactions has been the CH bond activation, activation of the coreactant CO2 or H2O over the support has not been considered an important issue. Additionally, of all investigated noble metals, Pt exhibits the highest activity for any cluster size used.44 Vines et al.47 emphasized the critical role of edge and corner sites of Pt nanoparticles in methane activation. Results of their DFT calculations as well as experimental observations showed that the usage of Pt nanoparticles (Pt79 clusters) instead of Pt surfaces (111) significantly facilitated the kinetics of methane dehydrogenation. They also explained this behavior by the presence of low-coordinated sites (edge, corner, and nearby sites) in Pt nanoparticles whose role has been to reduce the energy barriers of individual reaction steps and to stabilize the reaction intermediates. Vines et al.47 experimentally confirmed their theoretical predictions by investigating methane decomposition on Pt nanoparticles supported on an ordered CeO2 film on Cu (111) and by comparing their observations with previous experiments on Pt (111). It was shown that on supported Pt nanoparticles dehydrogenation of CH3 to CH and C species takes place already at 120 K while on the Pt (111) surface CH3 species remains stable up to 240 K. This indicates that on the Pt nanoparticles, the activation barrier for dehydrogenation has been significantly reduced. In summary, the prevailing understanding is that the rate limiting step in methane conversion is the activation of CH bond and the removal of the first H atom.48,49 The reactions in which only one or two of the CH bonds are broken have higher thermodynamic barriers than the reactions in which all four CH bonds of methane are cracked34 making it difficult to avoid complete oxidation of methane. Besides catalysis over transition metals and in homogeneous catalytic media, several alternative ways for methane activation have been tested such as photocatalytic conversion and dielectric barrier discharges which are discussed in more detail later. The question remains which is the most effective method to form CHx species, and then how to recombine these surface species into desired products without excessive energy requirements and carbon deposition.9

4. DRY REFORMING—CO2 AND CH4 TO SYNGAS Syngas, a mixture of hydrogen and carbon monoxide, is nowadays mainly produced by two major routes: partial oxidation (eq 1) of various hydrocarbons (natural gas, liquefied petroleum gas, naphtha, coal, petroleum coke, etc.) and steam reforming of methane (eq 2):50 Cn Hm þ ðn=2 þ m=4ÞO2 f nCO þ m=2H2 O CH4 þ H2 O f 3H2 þ CO ΔH or ¼ 206 kJ=mol

ΔGor ¼ 142 kJ=mol

ð1Þ ð2Þ

Suitable feedstocks are selected depending on the composition of the syngas required for downstream product to be produced (Tables 6 and 7). At industrial scale, syngas is used for the production of numerous chemicals, including methanol, dimethyl ether, FischerTropsch chemicals, ammonia, acetic 7093


Industrial & Engineering Chemistry Research Table 6. H/C Atomic Ratio of Typical Feedstock51 material

H/C

coke

0.13

charcoal

0.32

anthracite

0.38

bituminous coal

0.80

lignite

0.86

peat

1.15

heavy and residual oil

1.41

wood crude oil

1.44 1.71

lignite fuel oil

2.00

naphtha (light distillate feedstock)

2.18

liquefied petroleum gas (LPG)

2.67

liquefied natural gas (LNG)

3.43

methane

4.00

Table 7. Required H2/CO Ratio of Different Desired Products51 desired product

required H2/CO ratio

synthetic oil

1/22/1

methanol

2/1

acetic acid glycol

1/1 3/2

acetyloxide

1/1

propionic acid

4/3

methacrylic acid

5/4

ethanol

2/1

acetaldehyde

3/2

acetic ethene

5/4

acetic ether ethene

3/2 2/1

acid, and formic acid, which can potentially be used both as fuel and as feedstock for petrochemical industries. Noncatalytic partial combustion is an exothermic process, and it is performed at high pressure and high temperature. Because the necessary heat is generated by combustion of a portion of a feedstock, there is no need for external heat supply. However, both capital and operational costs are quite large. Due to the required very high reaction temperature (∼ 1770 K), special materials are needed as well as the purification of the produced syngas in order to remove soots and acid gases. In the resulting syngas, hydrogen to carbon monoxide ratio is lower. On the other hand, steam reforming is a highly endothermic process and it requires a large amount of heat input. A major drawback of the process is coke formation that occurs easily under the severe conditions at which the reaction is carried out: high pressure (2030 atm), high temperature >1170 K, and the lowest H2O/ CH4 ratio. Although using a large excess of water to methane could help reduce coking, it would decrease the thermal efficiency and negatively affect the economics of the process. As shown in Table 6, the H2/CO ratio in the syngas generated from methane is several times higher than that from coal. Accordingly, the main challenge to the research in this area is reducing the cost of production of syngas and reducing coke formation. For

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example, the syngas production cost is about 6070% of the capital cost of the methanol plant. Recent research directions under intensive investigation include catalytic partial oxidation of methane and CO2 reforming of methane.50 Reaction 3 below, between carbon dioxide and methane in production of syngas (that can be further used for making methanol 4 or Cn hydrocarbons 5) is called dry reforming. CH4 þ CO2 f 2H2 þ 2CO ΔH or ¼ 247 kJ=mol

ΔGor ¼ 170:8 kJ=mol

CO þ 2H2 f CH3 OH ΔH or ¼ 128:6 kJ=mol

ΔGor ¼ 29:1 kJ=mol ð4Þ

nCO þ 2nH2 f ðCH2 Þn þ nðH2 OÞ

ð3Þ

ð5Þ

This process received considerable attention since its successful realization would enable the direct conversion of natural gas into liquid fuels at remote extraction sites. Publications dealing with various aspects of this reaction are numerous.6,21,43,5257 Reaction 3 is limited by thermodynamics and even with the help of catalysts can take place only at high temperatures (> 1000 K).1 The first step of this process involves a highly endothermic reaction (ΔHr° = 247 kJ/mol) and the overall process may produce even more unwanted CO2 via the watergas shift reaction: CO þ H2 O f CO2 þ H2 ΔH or ¼ 41:1 kJ=mol


ð6Þ

Noble metals, such as Ru, Rh, Pd, Pt, and Ir, have been investigated extensively for the dry reforming reaction, but their high cost and limited availability hindered the commercialization of the process.56 Due to its fast turnover rates, long-term stability, and lower cost, nickel-based catalyst has been considered as the most suitable.6 However, the nickel catalyst is prone to deactivation by carbon deposition. To realize the process at an industrial scale, an effective catalyst that minimizes the coke deposition should be developed. It has been reported that the rate of coke formation, depending on the metal used, decreases in following order:10 Ni . Pt > Ru However, due to the high cost of Pt and Ru, a catalyst based on Ni would be more desirable. Recently NiCeZrO2 catalyst was reported to give high conversions of reactants (97%) and provide high resistance to catalyst deactivation but still requires a high temperature, 1073 K.1 Similarly to steam reforming, dry reforming suffers from the high energy requirements and inherent carbon deposition at such high temperatures, but it gives a lower ratio of H2/CO suitable for the FischerTropsch synthesis of long-chain hydrocarbons.54 By coupling the dry with steam (wet) reforming of methane, variable H2/CO ratios can be achieved. Furthermore, the large amounts of energy required could be supplied by coupling these reactions with the partial oxidation of methane, which is an exothermic process resulting in the socalled trireforming process.10 For example: CH4 þ 1=2O2 f 2H2 þ CO ΔH or ¼ 36 kJ=mol ΔGor ¼ 86:5 kJ=mol

ð7Þ

This combination of dry reforming, steam reforming, and partial oxidation of methane is considered a more promising and 7094



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attractive way of producing syngas compared to single reforming processes.

5. OXIDATIVE COUPLING OF CH4 WITH CO2 Oxidative coupling of methane (OCM) is also an attractive means of natural gas utilization. Lunsford58 gives a detailed review on the catalysts studied, possible mechanisms, and other important aspects of this reaction. The initial hydrocarbon produced by the reaction is the ethane, which is then converted in situ into the major target product, ethylene, and other higher hydrocarbons (C2þ products). Current ethylene annual production of around 140 million metric tons is based on steam cracking.59 Because this process is not only very energy intensive, but also emits 1.53 tons of CO2 per every ton of ethylene produced, alternative ways of ethylene production are of appealing interest for industry.59 OCM involves both surface and gas phase free-radical reactions.58 It is believed that the initial step— hydrogen abstraction from methane and methyl radical formation—requires the presence of surface oxygen.60 The ethane is then formed by gas phase coupling of CH3 3 radicals.4 However, these methyl radical intermediates are easily further oxidized by the gas-phase oxygen to CO2.60 A great challenge is finding the catalyst that would selectively break only one carbonhydrogen bond without burning it completely.59 Furthermore, developing other sources of oxygen that would provide surface oxygen without supplying the gas phase oxygen would consequently increase C2þ selectivity.60 Although achieved selectivity was high —around 80%—conversions were still very low for the commercialization of the process.59 On the other hand, the possibilities for nonoxidative methane coupling are reviewed in detail by Guczi et al.33 A suitable alternative to using oxygen could be carbon dioxide since it does not react with methyl radical in the gas phase.61 The coupling reactions between CH4 and CO2 can be then expressed as: 2CH4 þ CO2 f C2 H6 þ CO þ H2 O ΔH or ¼ 58:8 kJ=mol


ð8Þ

2CH4 þ 2CO2 f C2 H4 þ 2CO þ 2H2 O ΔH or ¼ 189:7 kJ=mol


ð9Þ

As reported for these reactions, thermodynamic calculations show that about 1535% of the equilibrium methane conversion can be achieved at 10731173 K at 1 atm and equimolar CH4/ CO2 mixture.61 By isotope tracer studies it has been shown that CO2 is converted to CO alone, and that C2þ hydrocarbons originated from CH4.60 Asami et al.61 did extensive studies over seventeen metal oxides (oxides of yttrium, lanthanum, samarium, titanium, zirconium, hafnium, niobium, chromium, manganese, iron, cobalt, copper, indium, aluminum, silicon, germanium, and bismuth). They concluded that CO2 and CH4 in most of the cases react overall with a ratio of around 2. Yttrium oxides gave the highest C2 yield (0.4%) and generally rare earth oxides showed high C2 selectivity (about 30%). Also, manganese, iron, cobalt, copper, and bismuth oxides are reduced during reaction while the other oxides remained unchanged.61Although methane activation by the active oxygen species supplied from the adsorbed carbon dioxide would definitely be an attractive route from the environmental perspective, further research is needed to bring this process to the industrial scale.

6. CATALYTIC PRODUCTION OF OXYGENATES Both homogeneous and heterogeneous catalytic systems have been investigated for possible direct conversion of methane and carbon dioxide to oxygenates. Most studies, however, have been devoted to the low-temperature production of acetic acid from methane and carbon dioxide. Currently, acetic acid is mainly produced from CH4 through the multistep process in which syngas and then methanol are produced first. The stoichiometry of the methanol carbonylation is simply the following: CH3 OH þ CO f CH3 COOH

ð10Þ

Prior to 1979, acetic acid was made using cobalt catalysts (BASF process) at 523 K and pressures up to 500700 atm. In 1970 Monsanto commercialized a rhodium carbonyl iodide catalyst which not only enabled significantly milder process conditions, around 450 K and 3040 atm, but also increased the selectivity toward acetic acid (>99%). The reaction mechanism has been studied in detail and it has been shown that besides transition metal component, HI is also necessary to catalyze the methanol conversion to methyl-iodide. Rh-based catalyst then catalyzes the carbonylation of methyl-iodide to acyl-iodide. The last step is the hydrolysis of the acyl-iodide in which HI is regenerated. 1

CH3 OH þ HI f CH3 I þ H2 O

2

CH3 I þ CO f CH3 COI

3

CH3 COI þ H2 O f CH3 COOH þ HI

Lately, this process has been replaced by the Cativa process62 which essentially follows the same mechanism except that it uses an iridium-based catalyst. Ir is more active than Rh, it also produces fewer byproducts and can operate at lower water levels, thereby reducing the overall costs. It is believed that the direct reaction of CH4 and CO2, if made possible, would be a much more attractive way to produce acetic acid, decreasing the production costs as well as environmental risks of the indirect route:63,64 CO2 þ CH4 T CH3 COOH ΔH °r ¼ 36:4 kJ=mol

ΔG°r ¼ 71:1 kJ=mol

ð11Þ

Such direct conversion of CH4 and CO2 to acetic acid is an example of maximum atomic efficiency. If successfully accomplished, it would eliminate the need for the intermediate step of syngas production. However, due to large positive values of the Gibbs free energy, the reaction is thermodynamically unfavorable at all conditions of practical interest. For example, for a 95% CO2 and 5% CH4 inlet mixture, equilibrium fractional conversion of CH4 at 1000 K and 100 atm is only 1.6 106.65 Effective ways of moving the equilibrium toward the production of acetic acid have to be found. Several research groups9,6470 have investigated different methods of acetic acid production by direct conversion of CH4 and CO2. Only a few have investigated the use of homogeneous catalysts. Kurioka and co-workers66 first reported the formation of acetic acid from CH4 and CO2 in Pd(OAc)2/Cu(OAc)2/ K2S2O8/CF3COOH catalytic system. Later Taniguchi et al.67 have found that this reaction can proceed in an aqueous solution 7095



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Table 8. Product Selectivities As a Result of Different Gas Feed9 CH4 and CO2 fed alternatively with

CH4 and CO2 fed alternatively without

CH4 and CO2 fed

H2 sweeping between the cycles

H2 sweeping between the cycles

simultaneously

48% C1C4 alcohols

36% formic acid

44% two cyclopentane derivatives

28% acetic acid

68% formaldehyde 21% cyclopentane derivative

5.6% acetic acid, etc.

21% methanol

12% acetic acid, etc.

5.5% methyl furan 3.4% cyclopentane derivative 2.4% butanone, etc.

of a vanadium catalyst, VO(acetylacetonate)2, in which also oxidant peroxydisulfate (K2S2O8) and trifluoroacetic acid (CF3COOH) solvent were used. The reaction was run in an autoclave at 353 K and both the catalyst and K2S2O8 had to be present for the reaction to occur. The authors reported very high yield of acetic acid based on methane (97%). However, they also noticed that the CO2 pressure does not have any effect on the yield and moreover that the reaction takes place even without the presence of CO2. This result motivated Wilcox et al.65,69 to question whether reaction 11 actually happens and to suggest a different scenario. As they claim, thermodynamically a more favorable reaction would actually be the reaction between methane and trifluoroacetic acid: CH4 þ CF3 COOH T CH3 COOH þ CF3 H

ð12Þ

Hence, their observations indicate that the solvent is not inert in this reaction system at the given experimental conditions.65 Moreover, since trifluoroacetic acid is expensive and difficult to handle, it would be desirable to find a more convenient solvent.7 Nizova et al.68 studied the carboxylation of methane with CO and CO2 in aqueous solutions containing O2 or H2O2, catalyzed by vanadium complexes. They reported the production of acetic acid and under certain conditions methanol, formaldehyde, etc. for both CO and CO2. However, resulting yields were low and despite mild temperatures (298373 K) pressure was high (around 50 atm). In U.S. patent 6960682,71 the authors claim the production of acetyl anhydrides as well as acetic acid from methane and carbon dioxide by contacting these two gases at 358 K in a high-pressure autoclave. Reaction took place in the presence of K2S2O8 as an initiator and VO(acac)2 as a catalyst, both of which were dissolved in an anhydrous acid and corresponding anhydride. These authors also mention British patent 226248GB by Dreyfus from 1924 in which the synthesis of acetic acid from methane and carbon dioxide over nickel carbonate as a catalyst was described. Furthermore, patent WO96/05163 of Hoechst A.G. was also cited in which the catalyst used contained metals from VIA, VIIA, and/or VIIIA group and high selectivities were claimed (7095%). Considering possible future industrial applications of production of acetic acid by direct conversion of methane and carbon dioxide, it would be more desirable to develop a heterogeneous catalytic system that would enable this reaction to give acceptable yields and selectivity. Several studies of direct low-temperature conversion of CH4CO2 on heterogeneous catalysts have been reported.9,6365 However, no detailed mechanism has been suggested, neither has the knowledge of what is actually happening on the surface of the catalyst significantly improved. Huang et al.9 applied periodic operation of the catalyst as a new and promising way to overcome thermodynamic limitations

in this reaction system. The catalyst used in their experiments was a coppercobalt based catalyst, otherwise used for hydrogenation of CO2. The choice of the catalyst was justified by the claim that the activation of methane produces the same intermediate adspecies (CHx, x = 0, 1, 2 or 3) as hydrogenation of CO2. Their experiments confirmed the formation of surface carbonaceous adspecies, CHx, during the reaction. Methane and carbon dioxide were introduced into the fixed bed reactor alternatively, under normal pressure and at isothermal conditions. Experiments were run at different Cu/Co ratios, from 0.3 to 6, and at different temperatures (373673 K). Different distribution of the products was observed depending on the way gases were introduced into the reactor (Table 8). Reported rates of formation of acetic and formic acids were 743 and 625 μmol/(gcat 3 h), respectively. The temperature of 523 K and Cu/Co ratio of 5 were the optimum conditions for the production of acetic acid. The same group continued their investigation by studying the reaction on a catalyst containing noble metals.64 The same stepwise approach was used, but this time on 1% and 2% Pd/ SiO2 and Rh/SiO2 catalysts. Experiments were run isothermally at different temperatures from 443 to 673 K. The authors consider the formation of MCH3 adspecies from methane and insertion of the CO2 molecule into the intermediate as two steps that limit the overall process (Figure 4). The results showed that the best catalyst among those investigated is 2% Pd/SiO2 at 473 K. It has been claimed that the good catalyst for this reaction should promote selective formation of intermediate MCH3 surface adspecies and also should insert CO2 readily. It has been shown that the loss of first H atom from the adsorbed CH4 on the surface of Pd is an exothermic process while the successive H loss is endothermic.64 On the other hand, on Rh surfaces successive losses of H atoms are all exothermic processes. Accordingly, the dominant intermediate specie on the Pd surface is desirable CH3 form, whereas on Rh surface besides CH3 there are also CH2 groups.64 Additionally, the adsorption energies of CHx species indicate that the insertion of CO2 molecule into the intermediate is more favorable on the Pd surface (Table 9).64 Again, quantitative results were presented for the formation rates of acetic acid which ranged from 2.5 to 38 μmol/(gcat 3 h). Huang et al.63 also tried to introduce oxygen into the reaction and, in that way, overcome thermodynamic limits: CH4 þ 0:5CO2 þ 0:5O2 f 0:75CH3 COOH þ 0:5H2 O ΔG°r ¼ 146:89 kJ=mol

ð13Þ

This time they used V2O5PdCl2/Al2O3 as a catalyst. Referring to the work of Taniguchi and Fujiwara, the authors claim that their system now contained the catalyst, the oxidant, and 7096



REVIEW

This reaction of acetic acid and acetylene to produce vinyl acetate 14 was tried experimentally by the same group:70 CH3 COOH þ C2 H2 f CH3 CO2 CHdCH2 ΔGor ¼ 65 kJ=mol

ð14Þ

which has been implemented in industry with Zn acetate supported on carbon as a catalyst. The resulting overall reaction with more favorable equilibrium would then be CO2 þ CH4 þ C2 H2 f CH3 CO2 CHdCH2 ΔGor ¼ 7 kJ=mol Figure 4. Schematic presentation of stepwise conversion of CH4 and CO2 to acetic acid.64

Table 9. Adsorption Energies (eV) of CHx Species with Transition Metals Ru, Rh, and Pd64 species

Ru

Rh

Pd 2.20

CH3

2.70

2.58

CH2

4.07

3.77

2.93

CH

5.64

5.45

4.46

C

5.51

5.47

4.47

the acidic medium. It was noticed that the catalyst activity increased with reaction time and the highest formation rates were around 30 μmol/(gcat 3 h) for given experimental conditions. Although these values are still low, no significant byproducts were observed. Wilcox et al.65 examined the synthesis of acetic acid by diffusereflectance FTIR when methane and carbon dioxide were introduced as equimolar mixture over 5% Pd/carbon and 5% Pt/Al2O3 catalysts. During the temperature programmed reaction (from 373 to 673 K), the estimated yield of acetic acid was ca. 1.5 106. Formation of acetic acid was observed only above 623 K for Pd/carbon catalyst and above 473 K for Pt/Al2O3 catalyst; no byproducts were recorded. Very low yields were reported and no mechanism for possible reaction pathways was suggested. Different pretreatment methods of catalyst (CO2 or CH4 flow) were explored and it was found that, in the case of Pd/ carbon catalyst, acetic acid was formed only when the catalyst was pretreated with CO2. However, various catalyst pretreatments did not exhibit any significant effect on the Pt/Al2O3 catalyst. To move the reaction equilibrium toward higher product formation, the authors suggested stepwise operation as it was done in the work of Huang et al.9 or coupling with another, thermodynamically more favorable reaction in which acetic acid is consumed. Thermodynamic analysis of possible coupling reactions performed using AspenPlus RGIBBS reactor model showed that the formation of methyl acetate from acetic acid and methanol was still limited (CH4 conversion of 1.81 103 at 300 K and 2000 atm) and although more favorable, the synthesis of acetic anhydride from ketene (CH4 conversion of 0.947 at only 300 K and 25 atm) would not be industrially economical.72 The synthesis of vinyl acetate from ethylene and oxygen would be a promising route (CH4 conversion of 0.99 at only 300 K and 1 atm) if a highly selective catalyst were developed that would suppress possible side reactions such as the oxidation of methane and ethylene.72

ð15Þ

As a result of their previous research, 5% Pt/Al2O3 was used as a catalyst as well as its mixture with Zn acetate/C. The authors offered the hypothesis that this reaction is sequential, namely, that in the first step acetic acid from methane and carbon dioxide has been produced and then in the second step vinyl acetate has been formed. Measured mol fractions of both acetic acid and vinyl acetate in the product gas were orders of magnitude higher that the equilibrium ones calculated at given conditions (T = 673 K and p = 1 atm) for an initially equimolar mixture of CO2, CH4, and C2H2.

7. DIELECTRIC BARRIER DISCHARGES (DBD) The group of Eliasson7377 raised attention to the application of quite an interesting technique of dielectric barrier discharges to the direct conversion of greenhouse gases to higher value products in order to circumvent thermodynamic barriers and facilitate the reaction progress. A schematic representation is given in Figure 5, and the method has been described in detail in Eliasson et al.73 Besides syngas, various distributions of products such as alkanes, alkenes, and oxygenates were examined as a function of different operating conditions and reactor configurations.7377As the authors claim, one of the main advantages of this reactor system is that reactions can be operated at low gas temperature and still achieve notably high methane conversion (up to around 68%). The power range of 200700 W was applied with a highvoltage generator running at around 30 kHz creating nonequilibrium plasma conditions. Conversions of CH4 and CO2 per kilowatt hour increased with the raise of the input power and the energy efficiency of the conversion increased 5 times.73 Measured temperature of the gas in the discharge zone as well as in the wall region (293423 K) showed no significant effect on the reactant conversions.73 The thermodynamic barrier has been overcome by formation of high-energy electrons (10 000100 000 K) that initiate the dissociation of reactant molecules. It has been assumed that the main mechanism is by free radical reactions. However, a more fundamental investigation is necessary because the reaction mechanism and kinetics are not yet fully understood. 8. PHOTOCATALYTIC CONVERSION One of the recent alternative approaches to overcoming thermodynamic limitations of some unfavorable reactions at mild conditions is the use of UV photo energy. As explained in great detail in Ravelli et al.,78 a photocatalyst is effective only in the excited state and the activation of the reagent involves either atom or electron transfer through the chemical transformation (Figure 6) on the lowest potential energy surface. 7097



Figure 5. Schematic representation of DBD reactor system.75

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achieve better conversions and product selectivities. Alternative routes—such as photocatalytic conversion and dielectric barrier discharges—require a detailed, overall energy balance assessment to determine their true competitiveness with traditional processes. Despite huge scientific interest as well as potential advantages of direct chemical utilization, an effective process still remains a great challenge. Further advances in understanding events on metal clusters on solid oxide surfaces via combined molecular modeling and experimental testing are needed.

’ AUTHOR INFORMATION Corresponding Author

*E-mail: [email protected]. Tel.: 314-935-8055.

’ REFERENCES

Figure 6. C acts as a photocatalyst in the tranformation RfP by absorbed light. An intermediate I is formed which gives I0 while regenerating C and forming end product P.78

The main advantages of photocatalytic processes are low energy consumption, stability of catalyst, safety, and stability of the reactor.1 A description of possibilities for photocatalytic conversion of methane can be found in Yuliati and Yoshida.1 One of the most widely used photocatalysts is titania because of its great (photo)chemical stability, low cost, and nontoxicity. Its relatively high electronhole recombination rate, which decreases its photoactivity, can be improved by doping with other metals or metal oxides which can act as electron traps and also can reduce the band gap energy shifting the absorption band to visible region.79 Shi et al.8 reported production of acetone from methane and CO2 over Cu/CdSTiO2/SiO2 catalyst. Despite significantly high selectivities (92.3% for acetone), conversion of methane was very low. Teramura et al.80 investigated the reduction of CO2 to CO in the presence of CH4 over MgO and ZrO2. It was reported that methane participated in the formation of surface acetate and formate. Resulting yields are still very low, which emphasizes the need for further improvement of the photocatalytic system. The major hurdle for industrial application is the variable and intermittent nature of the incident solar radiation. Despite the use of more advanced photoreactors designed for the maximization of light absorbance, which is critically important, lamp-driven photocatalysis is still more economically feasible than the use of solar light.78

9. SUMMARY Both reduction of carbon dioxide and low-temperature methane activation in the production of liquid fuels and high-value chemicals remain attractive targets for the “post petroleum” age. Up to now, the most successful conversion of methane and carbon dioxide has been the production of syngas (dry reforming), which is impaired by excessive carbon deposition and catalyst deactivation. Direct low-temperature production of acetic acid has been studied to a far lesser extent, and the proper understanding of the reaction pathways over catalytic surfaces is needed in order to

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Conversion of Methane and Carbon Dioxide to Higher Value Products

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