Evolution of Processes for Synthesis Gas Production - ACS Publications

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Evolution of Processes for Synthesis Gas Production: Recent Developments in an Old Technology Sebastia´ n C. Reyes,* John H. Sinfelt,† and Jennifer S. Feeley Corporate Strategic Research, ExxonMobil Research and Engineering Company, 1545 Route 22 East, Annandale, New Jersey 08801

The manufacture of gas mixtures of carbon monoxide and hydrogen has been a vitally important part of chemical technology for about a century. Originally, such mixtures were obtained by the reaction of steam with incandescent coke and were known as “water gas”. Used first as a fuel, water gas soon attracted attention as a source of hydrogen and carbon monoxide for the production of chemicals, at which time it gradually became known as synthesis gas. Eventually, steam reforming processes, in which steam is reacted with natural gas (methane) or a petroleum naphtha over a nickel catalyst, found wide application for the production of synthesis gas. A modified version of steam reforming known as autothermal reforming, which is a combination of partial oxidation near the reactor inlet with conventional steam reforming further along the reactor, improves the overall reactor efficiency and increases the flexibility of the process. Noncatalytic partial oxidation processes using oxygen instead of steam also found wide application for synthesis gas manufacture, with the special feature that they could utilize lowvalue feedstocks such as heavy petroleum residua. In recent years, catalytic partial oxidation employing very short reaction times (milliseconds) at high temperatures (850-1000 °C) is providing still another approach to synthesis gas manufacture, as Professor Lanny Schmidt and his students have shown in their pioneering exploratory research in this area. (See, for example, Hickman and Schmidt Science 1993, 259, 343.) Here, we consider some crucial issues in catalytic partial oxidation and report some new data with a bearing on these issues. Nearly complete conversion of methane, with close to 100% selectivity to H2 and CO, can be obtained with a Rh monolith under well-controlled conditions. Experiments on the catalytic partial oxidation of n-hexane conducted with added steam give much higher yields of H2 than can be obtained in experiments without steam, a result of much interest in obtaining hydrogen-rich streams for fuel cell applications. At 1000 °C, there is a negligible formation of carbon dioxide attributable to the overall reaction

Introduction The production of gas mixtures rich in carbon monoxide and hydrogen has been important industrially for a long time, roughly from about 1900 to the present. In an early method of obtaining such mixtures, a bed of hot coke was exposed alternately to blasts of air and steam.1 During the period of exposure to air, the burning of some of the coke heated the bed to a temperature in the vicinity of 1000 °C. After the air blast, the blast with steam brought about the overall reaction

C + H2O ) CO + H2

(1)

The use of steam in this step led to the early use of the term “water gas” in referring to the mixture of carbon monoxide and hydrogen formed. The reaction is endothermic (∆H°298 ) +31 kcal for graphitic carbon),2 in contrast to the highly exothermic reaction of carbon with oxygen that occurred during the air blast. The rise in temperature of the bed in the exothermic combustion step provided the heat for the subsequent endothermic reaction of the steam with coke. * Corresponding author. Phone: (908) 730-2533. E-mail: [email protected]. † Senior Scientific Advisor Emeritus.

C + 2H2O ) CO2 + 2H2

(2)

If equilibrium is assumed for both reactions 1 and 2, thermodynamic calculations of the composition of the gas mixture show that carbon dioxide and water are present in only trace amounts at 1000 °C, consistent with observations in the early water gas generators. Reaction 2, like reaction 1, is endothermic, but the heat of reaction is smaller (∆H°298 ) +21 kcal for graphitic carbon). We note that the equation for reaction 2 is a simple combination of the equation for reaction 1 and the equation for the so-called “water gas shift” reaction 3 of CO with H2O3

CO + H2O ) CO2 + H2

(3)

which is exothermic (∆H°298 ) -10 kcal). Because of their high heating values, water gas mixtures rich in carbon monoxide and hydrogen were originally of interest as fuels. Later, they attracted interest as sources of either H2 or CO, or of appropriate mixtures of the two, for the synthesis of various chemicals. A very important early example was the recovery of hydrogen from water gas for use in the catalytic synthesis of ammonia4 from elemental nitrogen

10.1021/ie0206913 CCC: $25.00 © 2003 American Chemical Society Published on Web 01/09/2003

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and hydrogen

N2 + 3H2 ) 2NH3

(4)

In this particular example, the nitrogen for the reaction could be obtained from the same operation as was used to generate the water gas mixture, because the gas issuing from the reactor during the air blast contained mostly nitrogen and carbon monoxide, the oxygen having been completely consumed. The combination of products formed during the alternate air and steam blasts of the coke thus consisted largely of a mixture of N2, CO, and H2, with some CO2. The concentration of CO in this mixture was then decreased to a very low level by exploitation of the water gas shift reaction. The equilibrium for this reaction was displaced in the direction of increased concentrations of CO2 and H2 by addition of excess H2O vapor to the mixture, coupled with cooling of the mixture to a temperature of about 500 °C. A satisfactory rate of reaction was obtained by contact of the mixture with an iron oxide catalyst. The concentration of CO was decreased to a level of about 1%, and the remaining CO, which is a severe poison for the metallic iron catalyst used in ammonia synthesis, was then removed by contacting the gas mixture with an ammoniacal solution of a cuprous salt.4-6 Prior to removal of the CO in this manner, CO2 was removed by contact with water under elevated pressure in an absorption tower. A later development for the removal of the last trace amounts of CO was its reaction with hydrogen to form methane over a nickel catalyst,7 as the methane is not a poison for the ammonia synthesis catalyst. An early example where both the hydrogen and the carbon monoxide from a water gas mixture were used in the production of an important chemical is the catalytic synthesis of methanol8

2H2 + CO ) CH3OH

(5)

using a mixture of zinc oxide and chromia as a catalyst. As in the case of ammonia synthesis, the reaction is exothermic and is conducted at high pressures for thermodynamic reasons. Other examples of important processes using both H2 and CO as reactants include the Fischer-Tropsch process for the production of hydrocarbons for fuels and lubricating oil base stocks as well as oxygenated hydrocarbons for chemicals,9,10 the “oxo” synthesis for the production of aldehydes in particular (hydroformylation),11 and the direct catalytic synthesis of higher alcohols using modified methanol synthesis catalysts.12 Because of the wide use of mixtures of H2 and CO for the synthesis of chemicals, it became common practice to refer to such mixtures as “synthesis gas” (or “syngas”).13,14 Steam Reforming In the 1920s and 1930s, the use of natural gas (methane) rather than incandescent coke in the production of synthesis gas began to attract attention.14,15 The overall reaction between methane and steam

CH4 + H2O ) CO + 3H2

(6)

is highly endothermic (∆H°298 ) +49 kcal). A catalyst, commonly containing nickel as the active component,

Figure 1. Water-gas shift and methane steam reforming equilibrium constants.

is used, and the process has come to be known as steam reforming. The process is typically operated with excess steam (H2O/CH4 mole ratio ) 2:1-4:1) at temperatures above about 800 °C. The composition of the product gas generally approaches that expected at equilibrium very closely. Both reactions 3 and 6 are important in determining the equilibrium composition. Figure 1 shows plots of the equilibrium constants as a function of temperature for these two important reactions. At a given reactor pressure and mole ratio of H2O to CH4 in the inlet gas stream to the reactor, the amount of carbon monoxide increases with increasing temperature, whereas the amounts of carbon dioxide and methane decrease. If the reactor pressure is increased, the amounts of methane and water increase substantially. Steam reforming was first applied commercially by the Standard Oil Company of New Jersey (now ExxonMobil) in 1930 at Baton Rouge, LA.16 In the 1950s and 1960s, interest began to develop in the use of light petroleum naphthas as feedstocks, particularly the highly paraffinic Middle East naphthas that were less desirable feedstocks for the production of aromatic hydrocarbons by another type of process known as catalytic reforming. The use of the same name, reforming, for the two very different kinds of processes can understandably be a source of confusion, but it should be clearly understood that they have no connection to one another. The catalytic reforming process is not operated with steam as a reactant or diluent. It is utilized for producing high-octane-number aromatic hydrocarbons for gasoline with the aid of so-called “bifunctional” precious metal catalysts.17-19 The steam reforming of light petroleum naphthas was originally of particular interest in regions of the world where natural gas was not readily available. Although steam reforming is used primarily for the production of hydrogen and carbon monoxide mixtures, a lowertemperature version of the process employing naphtha feedstocks attracted attention for the production of methane for so-called “town gas” for residential use. European cities were places where town gas manufacture was especially important. Dent and associates at the British Gas Council did considerable work in this area in the late 1950s.15 Highly active and stable nickel catalysts for application in town gas production were developed in the laboratories of the Exxon Research and Engineering Company in the 1960s.20-24 The use of these catalysts for hydrogen production from naphthas was also considered.25,26

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In steam reforming for town gas production, elevated pressures (30-40 atm) are employed and operating temperatures (400-450 °C) are much lower than those used for the production of synthesis gas. The molar ratios of steam to hydrocarbon are typically in the range of 8:1-12:1. The high ratios are important for limiting the formation of deactivating carbonaceous residues on the catalyst. If we consider n-hexane to be representative of the hydrocarbons in a light petroleum naphtha, we can think of the overall steam reforming process for producing methane as one that involves two major parts. The first part is the decomposition of the reactant hydrocarbon and the formation of CO and H2 as primary products. We represent this step by the overall reaction

which is highly endothermic (∆H°298 ) +228 kcal).2 The second part of the process consists of reaction 3 and the reverse of reaction 6, both of which are exothermic. As already mentioned, reactions 3 and 6 are close to being equilibrated in steam reforming operations, and the catalysts used in such operations have sometimes been called equilibration catalysts.27 Increasing the pressure and decreasing the temperature of the operation drive reaction 6 to the left, as desired for town gas applications. Steam reforming catalysts were early candidates for application in catalytic converters for decreasing the emissions of harmful pollutants in exhaust gases from automobiles.27 The chemistry in steam reforming is very versatile and important for many applications. Although steam reforming of petroleum naphthas declined in importance after huge quantities of natural gas were discovered in the North Sea area in the late 1960s, there has been a strong renewed interest in recent years in naphtha steam reforming for the production of hydrogen for fuel cells in connection with applications in electricpowered automobiles.28,29

of course, there is also some formation of CO2 and H2O. Reaction 8 is very exothermic (∆H°298 ) -118 kcal), contrasting markedly with the endothermicity of the steam reforming reaction of n-hexane, i.e., reaction 7. In the steam reforming section of an autothermal reformer, reactions 3 and 6 are equilibrated as in a conventional steam reformer, so that product compositions can be controlled by altering operations in a manner dictated by thermodynamic considerations. It is interesting that the basic idea of using an exothermic reaction to supply heat for an endothermic reaction is reminiscent of the alternate air and steam blasts in the old water gas generators. Another kind of process for the production of synthesis gas was also developed and commercialized in the 1940s by the Texaco32 and Shell33 Oil Companies. It is of particular interest for the gasification of low-quality feedstocks, such as petroleum residua and even coal or coke. Reaction with oxygen occurs in a refractory-lined vessel. There is commonly no addition of steam, and the reactor is not packed with a catalyst. Indeed, the operation is not conducive to the use of conventional catalysts. Oxygen rather than air is used to enable the reactors to reach temperatures high enough (about 1400 °C) to circumvent the problems with carbon formation. An O2/C ratio of about 0.7 is required for this purpose. As in the case of autothermal reforming, the design of the burners used in the process is a crucial issue. The reactors, which are commonly called gasifiers or thermal POX (partial oxidation) units, operate at pressures as high as 70 atm in the general range of temperature from 1200 to 1600 °C. The H2/CO ratios in the products are generally lower than about 2, and it might be necessary to make adjustments in the ratio by additional processing depending on the ultimate application of the synthesis gas. The technology is versatile and has been very widely employed since the initial applications. A very recent report indicates that several hundred gasifiers are in operation today.34

Autothermal Reforming and Partial Oxidation

Catalytic Partial Oxidation

From its very inception, steam reforming has been an important process, continually exhibiting advances in catalysts and in various engineering features of the reformers. One of these advances, autothermal reforming, was introduced by Haldor Topsøe in the late 1950s.30,31 It is a combination of partial oxidation and steam reforming, involving the introduction of both oxygen (or enriched air) and steam into a reforming reactor. The partial oxidation occurs in an inlet zone of the reactor, providing heat for the steam reforming reaction occurring in a second zone of the vessel that is packed with catalyst. Consequently, there is no need to supply heat to the reactor over and above the amount provided in the preheating of the reactants. This advance in steam reforming technology improves the overall reactor efficiency and increases the flexibility of the process. The ratio of molecules of oxygen gas to atoms of carbon in the hydrocarbon reactants (O2/C) is typically about 0.5-0.6, well below that required for complete combustion. The partial oxidation part of the process in the case of n-hexane as a feedstock is illustrated by the reaction

In recent years, the catalytic partial oxidation of hydrocarbons utilizing precious metals supported on porous ceramic monoliths has been the subject of much research. Professor Lanny Schmidt of the University of Minnesota has been a pioneer in this area.35-41 Other investigators are also now active in the field.42-46 Reaction times are very short (milliseconds), reminiscent of the old and widely used Ostwald process for the oxidation of ammonia on platinum or platinum-alloy gauzes in the commercial manufacture of nitric acid (ref 1, pp 597-598, and ref 13, pp 72-73). The monolith form of catalyst also resembles the kind of structural support used in catalytic converters for pollution control in automobile exhaust systems. In the case of methane oxidation, the desired reaction is

C6H14 + 6H2O ) 6CO + 13H2

C6H14 + 3O2 ) 6CO + 7H2

(7)

(8)

if the selectivity to CO and H2 were 100%. In practice,

CH4 + 0.5O2 ) CO + 2H2

(9)

which yields H2 and CO in a molar ratio of 2:1. Such a ratio is ideal for the reactant stream in methanol or Fischer-Tropsch synthesis. The direct partial oxidation of methane, in contrast to conventional steam reforming via reaction 6, is mildly exothermic (∆H°298 ) -8.5 kcal). Consequently, after the reactants are preheated to a desired inlet temperature, there is no need for additional input of heat through the reactor walls or


through provision for a separate inlet section of the reactor in which exothermic combustion reactions are conducted to provide heat for a downstream endothermic process such as reaction 6. Moreover, the direct catalytic partial oxidation reaction is much faster than the corresponding catalytic steam reforming reaction, by roughly 2 orders of magnitude.35-41 With an inlet reactant mixture of CH4 and O2 in a 2:1 molar ratio, thermodynamics indicates essentially complete methane conversion (>99%) to a 2:1 mixture of H2 and CO at temperatures above about 850 °C. However, the formation of H2 and CO as primary products does not proceed without any competition from the complete oxidation reaction to H2O and CO2. The competing reaction delays the attainment of the final equilibrium, with the result that some of the methane is still present in the product gases. The amount of unconverted methane is especially sensitive to the nature of the catalyst and the temperature to which the inlet reactant stream is preheated. Experiments have shown that the amount of unconverted methane decreases, and the selectivity to H2 and CO increases, when rhodium is used in place of platinum in the monolith. Similar positive effects are observed when the reactants are preheated to suitably high temperatures prior to their entrance into the reactor. The advantage of rhodium over platinum is attributed to a lower tendency of surface H atoms to be oxidized to surface hydroxyl radicals leading to the formation of water.47,48 As a result, desorption of the H atoms as H2 molecules becomes the favored process on rhodium. Work on catalytic partial oxidation employing monolith catalysts and very short reaction times at high temperatures is still in an exploratory stage, and prospects for industrial applications are not yet clear. However, there is much excitement in the area. In the next section of this paper, we present further discussion of important issues and report some new data with a bearing on these issues. Recent Results Concerning Key Issues in Catalytic Partial Oxidation In this section, we discuss some recent data on the catalytic partial oxidation of hydrocarbons using monolith catalysts. We focus on two examples that illustrate well some of the key factors influencing the product yields in high-temperature, short-contact-time reactors. In the first example, we examine the conversion of methane to synthesis gas. By making use of thermodynamic equilibrium calculations, we first identify conditions that favor high yields of H2 and CO with a molar ratio of 2. These mixtures are, of course, of great interest in the production of Fischer-Tropsch liquids and in the synthesis of methanol. In the second example, we study the generation of synthesis gas mixtures from n-hexane. This example is representative of an important class of reactions in which higher hydrocarbons are reacted with air and steam to produce a synthesis gas with a high hydrogen content. The use of steam as a co-reactant is beneficial for a number of reasons. It converts some of the carbon monoxide to carbon dioxide and additional hydrogen via the water gas shift reaction, it lowers the oxygen requirement (thereby decreasing the amount of nitrogen fed to the reactor), and it mitigates carbon formation on the catalyst surface. The generation of synthesis gas mixtures from higher hydrocarbons is now attracting considerable attention because, upon further

Figure 2. Equilibrium mole fractions of products for an inlet methane and oxygen mixture having an O2/C molar ratio of 0.5. Calculations with pure oxygen at a total pressure of 1 atm.

processing, a hydrogen-rich stream suitable for fuel cell applications can be produced. The processing usually entails additional water gas shift stages (with water addition) that drive the concentration of CO to very low levels, complemented by a final purification stage in which CO is substantially removed to levels that no longer interfere with the reactions on low-temperature fuel cell electrodes. The purpose of these two examples is to show that product compositions close to those at equilibrium can be obtained in very short reaction times, provided that suitable catalysts and conditions are used. The catalysts must be very active and very selective. In the transformation of methane to synthesis gas, for example, the disproportionate formation of H2O and CO2 through reactions in the gas phase or on the catalyst surface must be avoided because such reactions inevitably delay the attainment of equilibrium that dominantly favors H2 and CO as final products. In addition to the requirements of high catalyst activity and selectivity, another important requirement is a suitably high reaction temperature, which is paramount for achieving high yields of H2 and CO. In the partial oxidation of methane to synthesis gas, we would ideally like the transformation to proceed to completion via a direct path represented by the overall reaction 9. As shown by the results of thermodynamic calculations in Figure 2, H2 and CO are virtually the only products at equilibrium at very high temperatures. The equilibrium mole fractions in Figure 2 were calculated for an inlet reactant mixture of CH4 and O2 in a 2:1 molar ratio at atmospheric pressure. These mole fractions are shown as a function of temperature over the range 400-1000 °C. Consistent with expectations, the oxygen in the inlet reactant mixture is completely depleted. We emphasize the following points regarding the calculated mole fractions in Figure 2. First, the equilibria for both reactions 3 and 6 are completely taken into account. Second, equilibria involving carbon are not included in the calculations; i.e., equilibrium was not established for reactions such as 1 or 2, or for reactions 10 and 11 shown below

CH4 ) C + 2H2

(10)

2CO ) C + CO2

(11)

The justification for not including carbon is simply a

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Figure 3. Adiabatic outlet temperature as a function of inlet preheat temperature. Calculations are for an inlet methaneoxygen mixture having an O2/C molar ratio of 0.5 with either pure oxygen or air at a total pressure of 1 atm.

consequence of experience in catalytic partial oxidation and steam reforming operations, indicating that a steady state is approached in which the rate of formation of carbon is negligible compared to the rate of formation of the desired reaction products. The product compositions in such operations approach very closely the equilibrium compositions calculated without inclusion of carbon. As can be seen in Figure 2, as temperature increases, the conversion of methane increases, and the amounts of CO2 and H2O decrease. Such trends simply reflect the tendency of reaction 6 to convert CH4 and H2O and of reaction 3 in the reverse direction to consume CO2 as temperature increases. Although the highest yields of H2 and CO are realized at the highest temperatures, the matter of attaining such temperatures in a suitable manner requires some discussion. According to Figure 2, temperatures approaching 1000 °C are required to achieve essentially complete conversion to H2 and CO. Reaching this temperature level requires some provision for heat input because the desired overall reaction 9 is only mildly exothermic and therefore generates insufficient heat for the purpose. In Figure 3, the adiabatic outlet temperature is plotted as a function of the inlet preheat temperature for an inlet feed mixture having a CH4/O2 molar ratio of 2:1 with the use of pure oxygen in one case and air in the other. As the preheat temperature is increased from 300 to 600 °C, the outlet temperature increases from 805 to 920 °C for pure oxygen and from 710 to 820 °C for air. These calculated temperatures are a reflection of the lower amounts of H2O and CO2 present at equilibrium at higher temperatures and also of the diluting effect of nitrogen in the case of air. Thus, even in the ideal situation of an adiabatic operation, the results of Figure 3 indicate that high preheat temperatures are needed to bring the reaction temperature to a level of 1000 °C. These results also provide a reason for the use of pure oxygen instead of air. The previous discussion clearly illustrates the importance of reaching a sufficiently high reaction temperature. In our experience, the reactor outlet temperature is the critical parameter limiting the yields of H2 and CO from methane in laboratory experiments. However, there are several kinds of problems that contribute to undesirably low temperatures in such experiments. First, there is the difficulty of avoiding heat losses from

Figure 4. Schematic depiction of experimental monolith catalyst system employed in the present study.

small reactors. Laboratory studies typically employ cylindrical monoliths with diameters smaller than about 1 in. Second, there is the difficulty of premixing methane and oxygen rapidly and completely at high temperatures while avoiding gas-phase ignition. The complete oxidation products formed as a result of such ignition totally undermine the purpose of the experiment by eliminating the direct conversion of CH4 to H2 and CO in the subsequent action of the catalyst, thereby delaying the approach to the desired product equilibrium. Consequently, preheat temperatures in laboratory experiments have typically been lower than about 450 °C in previously reported studies. Also, if experiments are carried out with air instead of pure oxygen, as shown in Figure 3, the adiabatic temperature rise is substantially lower because of the added heat capacity of the nitrogen. In the present work, we have addressed the issue of reactor heat losses by employing a well-insulated quartz reactor within a three-zone furnace. We have also been able to preheat the reactants to high temperatures without ignition problems, and we have used pure oxygen instead of air. Figure 4 shows a schematic representation of the reactor used in our experiments. It corresponds to a standard setup in which the active monolith is sandwiched between two heat shields to preserve adiabaticity. The reactants are mixed and heated to about 150 °C prior to their entrance into the inlet zone of the reactor. Within the inlet zone of the reactor prior to contact with the catalyst, they are further preheated to the desired temperature by a controlled input of heat in the top section of the furnace. The temperature in each of the three zones within the furnace is independently controlled. For a rhodium catalyst supported on an alumina foam monolith, Figure 5 shows the methane conversion and selectivities to H2 and CO as a function of the preheat temperature. The preheat temperature is measured by the thermocouple located upstream of the front radiation shield. Figure 5 shows that very high yields of H2 and CO can be obtained at atmospheric pressure using CH4 and pure oxygen with a CH4/O2 molar ratio of 2:1


Figure 5. Experimental CH4 conversion and H2 and CO selectivities as a function of preheat temperature. Experiments carried out with an inlet methane and oxygen mixture having an O2/C molar ratio of 0.5. Pure oxygen, atmospheric pressure, GHSV ) 144 000 h-1.

Figure 6. Experimental outlet temperature as a function of preheat temperature. Experiments carried out with an inlet methane and oxygen mixture having an O2/C molar ratio of 0.5. Pure oxygen, atmospheric pressure, GHSV ) 144 000 h-1.

and a high gas hourly space velocity (GHSV ) 144 000 h-1). The GHSV is measured at standard conditions of pressure and temperature and is based on total monolith volume. Accounting for molar expansion and temperature rise we calculate that the contact time in these experiments is in the vicinity of 10 ms. The measured reactor outlet temperatures are given in Figure 6. The experimental results presented in Figures 5 and 6 are very consistent with the calculated values given in Figures 2 and 3. The conversion of methane increases as the reactor outlet temperature in the experiment increases. At a preheat temperature of 600 °C, the selectivities to H2 and CO are about 98 and 99%, respectively. The corresponding amounts of H2O and CO2 are very small but consistent with the values expected from equilibrium calculations. Because of the very high conversion of methane, yields of H2 and CO greater than about 95% are obtained. To our knowledge, this level of performance has not been reported earlier. Our experiments with methane and oxygen therefore demonstrate that very high yields of H2 and CO can be obtained at atmospheric pressure with foam monoliths using metallic rhodium as the active catalyst. One of the key remaining challenges in this area is the attainment of high yields at high pressures. This is necessary

Figure 7. Equilibrium mole fractions of products for an inlet n-hexane and air mixture having an O2/C molar ratio of 0.5 at a total pressure of 1 atm.

because the processes that utilize synthesis gas operate at high pressures. High pressures complicate the situation significantly. Because of the increase in the number of moles of species occurring in the production of synthesis gas from methane and other hydrocarbons, operations at higher pressure require higher temperatures to maintain yields of H2 and CO at a given level, as follows from Le Chaˆtelier’s principle. Higher temperatures, in turn, increase the potential for nonselective gas-phase reactions either during the mixing and preheating of the reactants or during the time in the reactor. Clearly, overcoming these difficulties requires the design of very specialized mixers that avoid gasphase ignition prior to entrance of the reactants into the catalyst zone and the development of very effective catalysts that permit the desired surface-catalyzed reactions to proceed more rapidly than the competing gas-phase reactions. The potential commercial application of this technology hinges heavily on these issues. We now turn to the example of the partial oxidation of n-hexane with air and steam to produce a synthesis gas rich in hydrogen. The analysis of this example has some similarities to that of the previous example on methane partial oxidation, but there are also some important differences. To understand better some of the features of this example, it is instructive to examine first the simpler situation where n-hexane and air are reacted without steam in a typical partial oxidation situation. Figure 7 shows the equilibrium mole fractions of products as a function of temperature for an inlet feed mixture having a C6H14/O2 molar ratio of 1:3 (i.e., O2/C ) 0.5) at a total pressure of 1 atm. This molar ratio corresponds to that of the reactants in the overall partial oxidation reaction 8. As in Figure 2, equilibria involving carbon are not included in the thermodynamic calculations. One of the first observations that can be made is that n-hexane is completely reacted to synthesis gas throughout the temperature range of interest for this type of reaction. The only hydrocarbon present in the equilibrium mixture is methane (below 800 °C), and this is a general feature in the partial oxidation of other hydrocarbon types as well. Thus, some of the conclusions drawn here for the reactions of n-hexane, air, and steam in monolith catalysts are applicable to the reactions of hydrocarbons in general. The nitrogen mole fraction is explicitly included in Figure 7 to remind us of nitrogen’s role in lowering the adiabatic temperature rise and in

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Figure 8. Equilibrium mole fractions of products for an inlet n-hexane, air, and steam mixture having an O2/C molar ratio of 0.5 and a H2O/C molar ratio of 1 at a total pressure of 1 atm.

lowering the partial pressures of the desired synthesisgas components. As in the previous example on methane, H2 and CO are the main products of the reactions of n-hexane with air at temperatures greater than about 800 °C. However, in the ideal situation described by reaction 8, the H2/CO molar ratio approaches a value of only 7/6 as a result of the lower H/C ratio in n-hexane compared to that in methane. The purpose of steam addition is therefore to exploit the water gas shift reaction 3 to produce additional hydrogen and carbon dioxide. Thermodynamic calculations can be used to determine how much steam should be added as a reactant for that purpose. Although the water gas shift reaction is exothermic, and thus not favored at high temperatures, increasing the steam partial pressure can increase the amount of CO transformed by the reaction. The steam that does not react, however, limits the reaction temperature by taking up some of the heat released by the oxidation reactions. These competing effects ultimately give rise to an optimal amount of steam that properly balances the extent of water gas shift conversion with the resulting reaction temperature. Figure 8 shows equilibrium mole fractions of products as a function of temperature for a typical situation in which steam is added to the reactant mixture. These compositions were calculated for an inlet mixture having a C6H14/O2 mole ratio of 1:3 (i.e., O2/C)0.5) and a C6H14/H2O mole ratio of 1:6 (i.e., H2O/C)1) at a total pressure of 1 atm. Contrasting the results in Figure 7 with those of Figure 8, one can clearly see the role of steam in producing additional H2 and CO2. The hydrogen mole fraction goes through a broad maximum when the temperature level reaches about 700 °C. This maximum in hydrogen occurs in the same temperature region where methane is no longer an equilibrium product. Thus, when steam is added to the reactant mixture, high yields of synthesis gas with a H2/CO molar ratio greater than about 2 can be obtained at a temperature of about 700 °C. Steam also helps reduce the amount of methane present at equilibrium. The drawback of this operation is that the mixture now contains CO2 and unreacted H2O, and depending on its ultimate use, some additional processing might be necessary. In the specific case of fuel cell applications, the final amounts of CO2 and H2O become even larger as a result of the lower-temperature watergas shift steps (with water addition) that are used to increase the H2 content further. Fortunately, the fuel

Figure 9. Adiabatic outlet temperature as a function of preheat temperature for an inlet n-hexane, air, and steam mixture with various O2/C and H2O/C molar ratios at a total pressure of 1 atm.

cell is largely unaffected by the presence of the N2, H2O, and CO2 that end up with the desired H2-rich stream. The critical issue with the fuel cell is the removal of CO, which is a poison for the catalyst surface in the anode. With regard to the control of the reaction temperature for optimal yields in these systems, there are some compensating effects that require discussion. They relate to reaction exothermicity and the requirements of air and steam. The key issues are illustrated in Figure 9 where the adiabatic outlet temperature is plotted as a function of the inlet preheat temperature. This figure shows results for three cases in which the O2/C and H2O/C ratios are varied within typical values. When an inlet mixture having an O2/C molar ratio of 0.5 and a H2O/C molar ratio of 1 is used, the adiabatic outlet temperatures exceed 800 °C even at preheat temperatures as low as 300 °C. This simply reflects the high exothermicity of the overall partial oxidation reaction 8, which is supplemented by additional heat release from the mildly exothermic water gas shift reaction. Because the results of Figure 8 suggest that temperatures in the neighborhood of 700 °C are needed for maximum yields of H2, one can consider lowering the amount of air or increasing the amount of steam. Lowering the air decreases the amount of nitrogen in the products, and additional steam increases the amount of hydrogen. Figure 9 shows that lowering the amount of air to decrease the O2/C molar ratio from 0.5 to 0.44 still leads to high enough reaction temperatures for optimum H2 yields. The simultaneous decrease of air to decrease the O2/C molar ratio from 0.5 to 0.44 and increase of steam from a H2O/C molar ratio of 1 to 2 decreases the temperature to levels that could become suboptimal at the lower preheat temperatures. It is emphasized that these trends are strictly valid under the ideal situation of adiabatic operation. In practice, any heat losses will tend to increase the amount of methane to the point where hydrogen yields might be compromised. It is advisable to aim for reaction temperatures slightly above the optimal values set by the theoretical analysis, especially when using small laboratory reactors where heat losses are a problem. Because of the broad maximum in hydrogen content of the product as a function of temperature (see Figure 8), the potential yield deficit arising from decreased watergas shift conversion at higher temperatures is small. This is consistent with the mild temperature depen-


Overall, the results of Figure 10 clearly demonstrate that hydrogen yields obtained in n-hexane partial oxidation with monolith catalysts containing rhodium as the active component improve substantially when water is added to the feed stream, as a result of the exploitation of the water gas shift reaction. Because of the dilution effect of both nitrogen and steam in the feed stream, gas-phase ignition is less of an issue than in operations without steam addition.

Figure 10. Experimental methane conversion and H2 and CO selectivities obtained in n-hexane experiments with air and steam at atmospheric pressure (upper field). Experimental values of O2/ C, H2O/C, and preheat temperatures used in n-hexane experiments with air and steam at atmospheric pressure (lower field).

dence of the equilibrium constant for the water gas shift reaction (see Figure 1). On the basis of these considerations, we carried out numerous experiments to study the reactions of nhexane with air and steam in monolith catalysts containing rhodium as the active component. The reactor setup was the same as that used in the methane partial oxidation experiments. As shown in the lower field of Figure 10, these experiments included systematic variations in the O2/C (0.4-0.54) and H2O/C (0-2.4) molar ratios, as well as in the preheat temperature (644-858 °C). A set of conditions (O2/C ) 0.44, H2O/C ) 1, T ) 700 °C, and GHSV ) 68 000 h-1) that was repeated several times during the sequence of experiments confirmed that the catalyst maintained activity with time on stream. The upper field of Figure 10 summarizes the trends in the selectivities to H2, CO, CO2, and CH4 for each of the conditions given in the lower field of the figure. For convenience, the hydrogen selectivity is defined as the amount of hydrogen in the products divided by the amount of hydrogen in the n-hexane reactant. A hydrogen selectivity in excess of 100% is therefore a measure of how much of the hydrogen in the steam is converted via the water gas shift reaction. In cases in which no steam is added as a reactant (i.e., under standard partial oxidation conditions), the hydrogen selectivity drops below 100% because some water is actually formed in the process (see Figure 7). One can also see that the amount of methane is generally small and consistent with the expected dependence on reaction temperature and on the amount of steam present in the reaction mixture. There is a very close correspondence between the pattern of variation of H2 selectivity in the upper field of Figure 10 and the pattern of changing H2O/C ratio in the lower field.

Concluding Remarks The production of synthesis gas mixtures very rich in carbon monoxide and hydrogen is aided by high temperatures. This has been clear since the earliest industrial operations employing water gas generators for the reactions of steam and air with red-hot coke. It is a consequence of thermodynamics. Thus, a mixture of steam and methane in a 1:1 molar ratio will yield an equilibrium mixture of hydrogen and carbon monoxide in a 3:1 molar ratio at 1000 °C, with only trace amounts of carbon dioxide, water, and methane being present. Similarly, a mixture of methane and oxygen in a molar ratio of 2:1 will yield a 2:1 equilibrium mixture of hydrogen and carbon monoxide at the same temperature. With a starting mixture of steam and methane, the high endothermicity of the steam-methane reaction requires that provision be made for the addition of heat well beyond the amount of preheat needed initially to bring the reactants to the desired temperature. Achieving this in the most efficient manner has been a key issue in the improvement of steam reforming operations over the years. With oxygen and methane as the starting mixture, the exothermicity of a methaneoxygen reaction removes this issue, the only consideration being the extent of reactant preheat required to supplement the heat released by the reaction in raising the temperature of the gas mixture to the desired level. The mechanism of the catalytic partial oxidation of methane to hydrogen and carbon monoxide at high temperatures is a matter of much interest.35-46 Consider the reaction of a mixture of methane and oxygen with a 2:1 molar ratio over a catalyst containing a precious metal (e.g., rhodium) as the active component. At temperatures much lower than those used in the production of synthesis gas, say in the range of 150200 °C, rhodium and other group VIII precious metals are very active oxidation catalysts, but only complete oxidation products (CO2 and H2O) are observed. What is happening at much higher temperatures with the monolith catalysts? One might simply argue that CO2 and H2O are still the initial reaction products, but that unconverted methane then enters into reaction with these products via the same chemistry that occurs in steam reforming, with the final product composition being governed by the equilibria for the water gas shift and methanation reactions. However, the reaction occurring in the monoliths is observed to be much faster (by roughly 2 orders of magnitude) than the steamreforming reaction on nickel catalysts, possibly implying that it takes a different course. Professor Lanny Schmidt and his students have concluded from their extensive work in this area that carbon monoxide and hydrogen are formed as primary reaction products.35-41 The rate of formation of these products relative to the rate of formation of complete oxidation products is therefore much higher than it is at low temperatures. The term direct partial oxidation is thus used in referring to this phenomenon.

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In a simple idealized situation, the selectivity of the catalyst for the direct partial oxidation reaction would be so high that conversion of the methane would proceed completely to hydrogen and carbon monoxide without any significant occurrence of other reactions. The extent to which this simple situation is approached in a given case has not been clearly established. Although very high selectivities to H2 and CO have been obtained with nearly complete conversion of methane, the absence of extensive data at low to moderate conversion levels precludes one from drawing firm conclusions about the possible contributions of other reactions as the rapid conversion of the methane proceeds. The possibility of complete oxidation products contributing significantly to the ultimate formation of H2 and CO is not readily dismissed. However, regardless of the present state of understanding of this matter, the approach of using monolith catalysts and millisecond reaction times at high temperatures has been a major development in the catalytic partial oxidation of hydrocarbons. As demonstrated by results reported in this paper, at a methane conversion close to 97%, selectivities to H2 and CO of 98-99% have been obtained with a rhodium monolith catalyst at atmospheric pressure. This has been achieved by careful attention to experimental details such as preheating the reactants to a suitably high temperature to ensure that the temperature reached in the reactor is high enough for essentially complete conversion of the methane to hydrogen and carbon monoxide. The minimization of reactor heat losses from the small laboratory reactors used in obtaining the results, coupled with still other experimental measures pointed out in the discussion of the results, has been crucial for obtaining methane oxidation data under well-controlled conditions. In recent years, there has been much interest in obtaining hydrogen from hydrocarbons in the gasoline boiling range for application in fuel cells for powering automobiles. In addition to the results reported in this paper, some results of Schmidt in this regard have been reported recently.49,50 This approach would have the advantage of exploiting the existing infrastructure for dispensing gasoline to the consumer. For an “on-board” application in an automobile, high-temperature catalytic partial oxidation could find application here. A variation of the latter in which steam is added along with air to the hydrocarbon fuel stream admitted to the reactor provides a way to increase substantially the amount of hydrogen obtained in such an operation. The added water is clearly the source of the increased amount of hydrogen, as the total production of hydrogen is much higher than what can be obtained from the hydrocarbon fuel alone. This has been demonstrated very clearly by the results presented in this paper. The participation of the added water in the overall process is similar to that already mentioned in connection with a possible role of complete oxidation products in secondary reactions occurring in the basic version of the process operated without the addition of water. Thus, we see advantages again for utilizing both air and steam in a reactor for the production of synthesis gas. This is a theme that has surfaced repeatedly in a number of processes that differ in detail but depend in an important way on the use of both of these common reagents. The theme goes all the way back to the water gas generators used in the early 1900s.

In closing this discussion of synthesis gas production, we emphasize the decisive role that it has played in chemical technology for the past century and that it appears destined to play in major technological innovations envisioned for the future. The new results presented in this paper provide a flavor for the kind of research currently done in connection with exciting new applications of old “water gas” chemistry. Acknowledgment The authors thank Mrs. Carol Bordok and Mr. Bruce A. DeRites for the generation of the experimental data on methane and n-hexane partial oxidation, respectively, presented in this paper. Many thanks also go to Dr. Paul J. Berlowitz for helpful discussions and guidance in the n-hexane study, which is part of a broader set of activities in the production of hydrogen-rich streams from hydrocarbons for fuel cell applications. Literature Cited (1) Briscoe, H. T. General Chemistry for Colleges, 3rd ed.; The Riverside Press: Cambridge, MA, 1943: pp 138, 139. (2) Glasstone, S. Thermodynamics for Chemists; D. van Nostrand Company, Inc.: New York, 1947: pp 73, 77. (3) Rideal, E. Concepts in Catalysis; Academic Press: New York, 1968: pp 177, 178. (4) Frankenburg, W. G. The Catalytic Synthesis of Ammonia from Nitrogen and Hydrogen. In Catalysis; Emmett, P. H., Ed.; Reinhold Publishing Corporation: New York, 1955; Vol. 3, pp 171263. (5) Gould, E. S. Inorganic Reactions and Structure; Henry Holt and Company: New York, 1955: p 158. (6) Cotton, F. A.; Wilkinson, G. Advanced Inorganic Chemistry, 1st ed.; John Wiley and Sons: New York, 1962: p 752. (7) Taylor, H. S. Industrial Hydrogen; The Chemical Catalog Company, Inc.: New York, 1921: pp 178-188. (8) Natta, G. Synthesis of Methanol. In Catalysis; Emmett, P. H., Ed.; Reinhold Publishing Corporation: New York, 1955; Vol. 3, pp 349-411. (9) Berkman, S.; Morrell, J. C.; Egloff, G. Catalysis; Reinhold Publishing Corporation: New York, 1940: pp 1042-1044. (10) Storch, H. H.; Golumbic, N.; Anderson, R. B. The FischerTropsch and Related Syntheses; John Wiley and Sons: New York, 1951. (11) Wender, I.; Sternberg, H. W.; Orchin, M. The Oxo Reaction. In Catalysis; Emmett, P. H., Ed.; Reinhold Publishing Corporation: New York, 1957; Vol. 4, pp 73-130. (12) Natta, G.; Colombo, U.; Pasquon, I. Direct Catalytic Synthesis of Higher Alcohols from Carbon Monoxide and Hydrogen. In Catalysis; Emmet P. H., Ed.; Reinhold Publishing Corporation: New York, 1957; Vol. 4, pp 131-174. (13) Prettre, M. Catalysis and Catalysts; Dover Publications: New York, 1963; pp 73, 74 (translation by David Antin of the third edition of Catalyse et Catalyseurs, originally published in 1946 by Presses Universitaires de France, Paris, France). (14) Storch, H. H. Synthesis Gas from Methane, Oxygen, and Steam. In The Chemistry of Petroleum Hydrocarbons; Brooks, B. T., Boord, C. E., Kurtz, S. S., Jr., Schmerling, L., Eds.; Reinhold Publishing Corporation: New York, 1955; pp 357-364. (15) Rostrup-Nielsen, J. R. Catalytic Steam Reforming. In Catalysis, Science and Technology; Anderson, J. R., Boudart, M., Eds.; Springer-Verlag: Berlin, 1984; Vol. 5, pp 1-117. (16) Byrne, P. J., Jr.; Gohr, E. J.; Haslam, R. T. Recent Progress in Hydrogenation of Petroleum. Ind. Eng. Chem. 1932, 24, 1129. (17) Sinfelt, J. H. Bifunctional Catalysis. In Advances in Chemical Engineering; Drew, T. B., Hoopes, J. W., Jr., Vermeulen, T., Eds.; Academic Press: New York, 1964; Vol. 5, pp 37-74. (18) Sinfelt, J. H. Catalytic Reforming of Hydrocarbons. In Catalysis, Science and Technology; Anderson, J. R., Boudart, M., Eds.; Springer-Verlag: Berlin, 1981; Vol. 1, pp 257-300. (19) Sinfelt, J. H. Catalytic Reforming. In Handbook of Heterogeneous Catalysis; Ertl, G., Kno¨zinger, H., Weitkamp, J., Eds.; Wiley-VCH Verlagsgesellschaft mbH: Weinheim, Germany, 1997; Vol. 4, pp 1939-1955.

Ind. Eng. Chem. Res., Vol. 42, No. 8, 2003 1597 (20) Taylor, W. F.; Sinfelt, J. H. High Activity Nickel-Alumina Catalyst. U.S. Patent 3,320,182, 1967. (21) Taylor, W. F.; Sinfelt, J. H. Promoted Catalyst for Methane Production. U.S. Patent 3,407,149, 1968. (22) Taylor, W. F.; Sinfelt, J. H. Methane Production Using a Promoted Catalyst. U.S. Patent 3,404,100, 1968. (23) Taylor, W. F.; Sinfelt, J. H. Promoted Catalyst Used for Town Gas Production. U.S. Patent 3,395,104, 1968. (24) Sinfelt, J. H.; Taylor, W. F. Controlled Conversion of Light Naphtha to Town Gas. U.S. Patent 3,450,514, 1969. (25) Taylor, W. F.; Sinfelt, J. H. Selective Partial Conversion of Naphtha Hydrocarbons to Hydrogen. U.S. Patent 3,394,086, 1968. (26) Taylor, W. F.; Sinfelt, J. H. Catalyst for Low-Temperature Conversion of Hydrocarbons to Hydrogen and Methane. U.S. Patent 3,396,124, 1968. (27) Taylor, W. F.; Sinfelt, J. H. Equilibration and Steam Reforming Catalysts. U.S. Patent 3,444,099, 1969. (28) Berlowitz, P. J.; Darnell, C. P. Fuel Choices for Fuel Cell Powered Vehicles; SAE Technical Paper 2000-01-003; SAE International: Warrendale, PA, 2000. (29) Pettersson, L. J.; Westerholm, R. State of the Art of MultiFuel Reformers for Fuel Cell Vehicles: Problem Identification and Research Needs. Int. J. Hydrogen Energy 2001, 26, 243. (30) Rostrup-Nielsen, J. R. Production of Synthesis Gas. Catal. Today 1993, 18, 305. (31) Christensen, T. S.; Primdahl, I. I. Improve Syngas Production Using Autothermal Reforming. Hydrocarbon Process. 1994, 73 (Mar), 39-46. (32) Falsetti, J. S. Gasification Process for Maximizing Refinery Profitability. Hydrogen Technol. Int 1993, 57-61. (33) Eilers, J.; Posthuma, S. A.; Sie, S. T. The Shell Middle Distillate Synthesis Process (SMDS). Catal. Lett. 1990, 7, 253. (34) Simbeck, D. Gasification Technology Update. Paper presented at the Fifth IChemE European Gasification Conference, Noordwijk, The Netherlands, Apr 8-10, 2002. (35) Hickman, D. A.; Schmidt, L. D. Production of Syngas by Direct Catalytic Partial Oxidation of Methane. Science 1993, 259, 343. (36) Hickman, D. A.; Schmidt, L. D. Steps in CH4 Oxidation on Pt and Rh Surfaces: High-Temperature Reactor Simulations. AIChE J. 1993, 39, 1164. (37) Hickman, D. A.; Schmidt, L. D. Synthesis Gas Formation by Direct Oxidation of Methane over Pt Monoliths. J. Catal. 1992, 138, 267. (38) Hickman, D. A.; Haupfear, E. A.; Schmidt, L. D. Synthesis Gas Formation by Direct Oxidation of Methane over Rh Monoliths. Catal. Lett. 1993, 17, 223.

(39) Deutschmann, O.; Schmidt, L. D. Modeling the Partial Oxidation of Methane in a Short Contact Time Reactor. AIChE J. 1998, 44, 2465. (40) Schmidt, L. D.; Deustchmann, O.; Goralski, C. T., Jr. Modeling the Partial Oxidation of Methane to Syngas at Millisecond Contact Times. Stud. Surf. Sci. Catal. 1998, 119, 685692. (41) Deutschmann, O.; Schmidt, L. D. Two-Dimensional Modeling of Partial Oxidation of Methane on Rhodium in a Short Contact Time Reactor. In Twenty-Seventh Symposium (International) on Combustion; The Combustion Institute: Pittsburgh, PA, 1998; pp 2283-2291. (42) Buyesvskaya, O. V.; Wolf, D.; Baerns, M. RhodiumCatalyzed Partial Oxidation of Methane to CO and H2. Transient Studies on Its Mechanism. Catal. Lett. 1994, 29, 249. (43) Walter, K.; Buyesvskaya, O. V.; Wolf, D.; Baerns, M. Rhodium-Catalyzed Partial Oxidation of Methane to CO and H2. In Situ DRIFTS Studies on Surface Intermediates. Catal. Lett. 1994, 29, 261. (44) Ostberg, M.; Aasberg-Petersen, K.; Basini, L.; Guarinoni, A. Catalytic Partial Oxidation of Natural Gas. Presented at the 2002 AIChE Spring Meeting, New Orleans, LA, Mar 10-14, 2002. (45) Mallens, E. P. J.; Hoebink, J. H. B. J.; Marin, G. B. The Reaction Mechanism of the Partial Oxidation of Methane to Synthesis Gas: A Transient Kinetic Study over Rhodium and a Comparison with Platinum. J. Catal. 1997, 167, 43. (46) Hofstad, K. H.; Hoebink, J. H. B. J.; Holmen, A.; Marin, G. B. Partial Oxidation of Methane to Synthesis Gas over Rhodium Catalysts. Catal. Today 1998, 40, 157. (47) Williams, W. R.; Marks, C. M.; Schmidt, L. D. Steps in the Reaction H2 + O2 T H2O on Pt: OH Desorption at High Temperatures. J. Phys. Chem. 1992, 96, 5922. (48) Zum Mallen, M. P.; Williams, W. R.; Schmidt, L. D. Steps in H2 Oxidation on Rh: OH Desorption at High Temperatures. J. Phys. Chem. 1993, 97, 625. (49) Schmidt, L. D. Chemicals from Alkanes in Millisecond Reactors. Presented at the 17th North American Catalysis Society Meeting, Toronto, Ontario, Canada, Jun 3-8, 2001. (50) Schmidt, L. D. University of Minnesota, Minneapolis, MN. Private communication, Oct 2002.

Received for review September 9, 2002 Accepted October 29, 2002 IE0206913

Evolution of Processes for Synthesis Gas Production - ACS Publications

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