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Pamela Spath and David Dayton

In principle, syngas can be produced from any hydrocarbon feedstock, including: natural gas, naphtha, residual oil, petroleum coke, coal, and biomass. The lowest cost routes for syngas production, however, are based on natural gas, the cheapest option being remote or stranded reserves. Economic considerations dictate that the production of liquid fuels from syngas today translates into the use of natural gas as the hydrocarbon source. Nevertheless, the syngas production operation in a gas-to-liquids plant amounts to greater than half of the capital cost of the plant. The choice of technology for syngas production also depends on the scale of the synthesis operation. Syngas production from solid fuels can require and even greater capital investment with the addition of feedstock handling and more complex syngas purification operations. The greatest impact on improving the economics of gas-to liquids plants can be had by decreasing capital costs associated with syngas production and improving the thermal efficiency with better heat integration and utilization. Improved thermal efficiency can be obtained by combining the gas-to-liquids plant with a power generation plant to take advantage of the availability of low-pressure steam. The extensive research and development efforts devoted to syngas conversion to fuels and chemicals are documented in a vast amount of literature that tracks the scientific and technological advancements in syngas chemistry. The purpose of this report is to review the many syngas to products processes and summarize the salient points regarding the technology status and description, chemistry, catalysts, reactors, gas cleanliness requirements, process and environmental performances, and economics. Table 1 lists the products examined in this study and gives some facts about the technology as well as advantages and disadvantages. Table 2 summarizes the catalysts, process conditions, conversions, and selectivities for the various syngas to products processes.

Table 1: Summary of Product Information Facts, Advantages, & Disadvantages Product


- Largest use of syngas - Predominately made via SMR

- Appears to be the most cost competitive option for biomass - Automakers working on H2 fueled vehicles - H2 itself is a clean fuel - Concentrated CO2 stream - Mature technology - High yields - Yields petroleum products consistent with todays fuel supply - Successful operation of slurry phase reactor - No sulfur and low aromatic products - Reduced tailpipe emissions compared to conventional diesel - Mature technology - Mature technology - Many uses for MeOH (formaldehyde, acetic acid, DME, MTBE, MTG, MTO, MOGD) - Fuel used as M100 and M85 - Possible source of H2 & MeOH reforming is a practiced technology - In situ dehydration of MeOH to DME increases MeOH yields - Many potential uses (diesel or cooking fuel, refrigerant, chemical feedstock)

- Delivered price can be significantly greater than plant gate price



- Existing plants in South Africa, Malaysia, and New Zealand - Predominately made from coal but South Africa switching some operating plants to natural gas - Many companies are examining FTL for stranded natural gas - Largest consumer of H2 - 2nd largest synthetic chemical product - Predominately from syngas made via SMR - Several automakers have developed prototype direct MeOH fuel cell vehicles - Commodity chemical

- Low yields - Unavoidable broad range of products; poor selectivity

- Toxic - Have to compete with mega methanol plants - LPMeOH still in development stage - Poor solubility in gasoline and phase separation problems - M85 has not developed on a large scale



- Largest consumer of MeOH - Demand driven by construction industry

Acetic acid MTO and MOGD MTG

- Half of worldwide production comes from MeOH carbonylation

- Low aromatic content in MOGD gasoline - Selectivity of olefins to gasoline & diesel > 95% for MOGD - High quality, high octane gasoline

- No commercial MTO or MOGD plants but UOP and HYDRO does license a process that produces ethylene and propylene - High amount of aromatics in gasoline - High decalene concentrations - No commercial plants - TIGAS yields a lower quality gasoline than MTG - No commercial plants from syngas - Many universities have worked on fermentor designs but present status of work for most is unknown No commercial plants Low yields Poor selectivity Currently, there is little activity in mixed alcohol synthesis - Very versatile process -

- One plant in New Zealand built in 1985 but currently it produces only MeOH - MeOH to gasoline - Currently produced via direct fermentation of carbohydrates or hydration of ethylene - Alcohol fuels developed by many - Some tested in Europe

TIGAS Ethanol

Mixed alcohols

- Is currently a fuel additive - Currently, has a tax credit - Compared to direct fermentation, able to process nearly any biomass resource - More attractive blending stock than MeOH - Lower CO:H2 ratio required - Lower CO:H2 ratio required

Oxosynthesis products

Isosynthesis products

- Hydroformylation of olefins with syngas - Produces C3-C15 aldehydes - Aldehydes typically converted to alcohols and acids - Converts syngas to isobutene and isobutane

- Lower CO:H2 ratio required

- No commercial plants

Table 2: Summary of Syngas Conversion Processes and Conditions

Process Conditions
Process Catalyst Fe Fischer-Tropsch Synthesis Co Ru ZnO/Cr2O3 Cu/ZnO/Al2O3 Fe/FeO + additives Alkali/ZnO/Cr2O3 Alkali/Cu/ZnO(Al2O3) Alkali/CuO/CoO Alkali/MoS2 Co carbonyl Co P modified Oxosynthesis Rh P modified Isosynthesis Steam Methane Reforming ThO2 ZrO2 Ni 60-120 400-450 300-425 850 7-25 100-1000 (300) 350 15-30 1:1 + propylene 0.85:1 1:1 na 40-50 i-C4 30 100% CH4 conversion 15 Syngas/ hydrogen T (C) 300-350 200-240 350 220-275 430-480 (550 max) 300-425 275-310 260-340 260-350 110-200 160-200 P (bar) 10-40 7-12 250-350 50-100 100-500 125-300 50-100 60-200 30-175 200-300 50-100 H2:CO 1.7:1 2.15:1 3:1 99% (25% max/pass 4-7% actual/pass) 10-35/pass 5-20% 20-30 5-30 10 % conv (CO basis) 50-90 with recycle Product -olefins gasoline Waxes diesel Waxes Methanol ammonia Branched primary alcohols Primary alcohols Linear primary alcohols linear alcohols C11-C14 alcohols C4 aldehydes > 99% with recycle Selectivity ASF -48% (max) 15-40% actual ASF - 40% (max) > 99% with recycle

Methanol Synthesis Ammonia

2-3:1 H2:N2 1:1 2-3:1 0.5-4:1 1:1 1:1 + olefin 1:1 + olefin

Higher Alcohol Synthesis

30-45% C2+ 17-25% CO2 ASF 75-90% C2+ in liquid product

> 90%

In its simplest form, syngas is composed of two diatomic molecules, CO and H2, that provide the building blocks upon which an entire field of fuel science and technology is based. Over the years, the gaseous mixture of CO and H2 has had many names depending on how it was formed; producer gas, town gas, blue water gas, synthesis gas, and syngas, to name a few. In the 1800s coal gasification was used to provide much of the syngas used for lighting and heating. The beginning of the 20th century saw the dawn of fuels and chemicals synthesis from syngas. The synthesis of hydrocarbons from CO hydrogenation was discovered in 1902 by Sabatier and Sanderens who produced methane by passing CO and H2 over Ni, Fe, and Co catalysts. At about the same time, the first commercial hydrogen from syngas produced from steam methane reforming was commercialized. Haber and Bosch discovered the synthesis of ammonia from H2 and N2 in 1910 and the first industrial ammonia synthesis plant was commissioned in 1913. The production of liquid hydrocarbons and oxygenates from syngas conversion over iron catalysts was discovered in 1923 by Fischer and Tropsch. Variations on this synthesis pathway were soon to follow for the selective production of methanol, mixed alcohols, and isosynthesis products. Another outgrowth of Fischer-Tropsch Synthesis (FTS) was the hydroformylation of olefins discovered in 1938. Much of the syngas conversion processes were being developed in Germany during the first and second world wars at a time when natural resources were becoming scare and alternative routes for hydrogen production, ammonia synthesis, and transportation fuels were a necessity. Thus the poor economics of some of these syngas conversion processes were of little consequence. With the development of the petroleum industry in the 1940s and beyond, the unattractive economics of many of these syngas routes became an issue and were replaced by petroleum-based processes. To this day, however, methanol and ammonia are still produced from syngas using essentially the same processes originally developed and, apart from hydrogen production, constitute the major uses of syngas. While petroleum-derived transportation fuels and chemicals were taking center stage in the worldwide market, certain political and environmental drivers were helping to develop improved commercial syngas conversion processes. The political situation in South Africa and the abundance of local coal reserves in that country helped bring about the most successful commercial FTS industry in the world based on syngas production from coal gasification. Sasol currently supplies diesel, gasoline, and other high-value hydrocarbons to local and global markets. The Arab oil embargo in the 1970s also highlighted the US dependence on foreign oil imports and helped renew interest in syngas conversion technologies. The increasing environmental concerns and tighter regulations surrounding fossil fuel use also provided impetus for syngas conversion technologies to produce cleaner (virtually no sulfur) fuels and chemicals. The use of methanol and isobutene for the production of methyl tert-butyl ether (MTBE), an octane enhancing oxygenated component in reformulated gasoline, also increased demand for syngas conversion technologies. MTBE itself, however, is becoming an environmental concern as a

watershed pollutant and the future use of this oxygenate is uncertain. The latest environmental driver to likely increase demand for syngas even more is the goal of establishing a hydrogen economy. The vision is for hydrogen to be the fuel of choice for transportation and electricity generation via high efficiency, environmentally benign fuel cells. In principle, syngas can be produced from any hydrocarbon feedstock, including: natural gas, naphtha, residual oil, petroleum coke, coal, and biomass. The lowest cost routes for syngas production, however, are based on natural gas, the cheapest option being remote or stranded reserves. Economic considerations dictate that the production of liquid fuels from syngas today translates into the use of natural gas as the hydrocarbon source. Nevertheless, the syngas production operation in a gas-to-liquids plant amounts to greater than half of the capital cost of the plant. The choice of technology for syngas production also depends on the scale of the synthesis operation. Syngas production from solid fuels can require and even greater capital investment with the addition of feedstock handling and more complex syngas purification operations. The greatest impact on improving the economics of gas-to liquids plants can be had by decreasing capital costs associated with syngas production and improving the thermal efficiency with better heat integration and utilization. Improved thermal efficiency can be obtained by combining the gas-to-liquids plant with a power generation plant to take advantage of the availability of low-pressure steam. The syngas composition, most importantly the H2/CO ratio, varies as a function of production technology and feedstock. Steam methane reforming yields H2/CO ratios of 3/1 while coal gasification yields ratios closer to unity or lower. Conversely, the required properties of the syngas are a function of the synthesis process. Fewer moles of product almost always occur when H2 and CO are converted to fuels and chemicals. Consequently, syngas conversion processes are more thermodynamically favorable at higher H2 and CO partial pressures. The optimum pressures depend on the specific synthesis process. With the exception of methane steam reforming, catalytic syngas conversion processes are exothermic reactions generating large excesses of heat. This highlights the specific need for removing this heat of reaction to carefully control reaction temperatures to maintain optimized process conditions. Maximizing product yields, minimizing side or competing reactions, and maintaining catalyst integrity dictate optimum synthesis reaction temperatures. Catalysts play a pivotal role in syngas conversion reactions, in fact, fuels and chemicals synthesis from syngas does not occur in the absence of appropriate catalysts. The basic concept of a catalytic reaction is that reactants adsorb onto the catalyst surface and rearrange and combine into products that desorb from the surface. One of the fundamental functional differences between syngas synthesis catalysts is whether or not the adsorbed CO molecule dissociates on the catalyst surface. For FTS and higher alcohol synthesis, CO dissociation is a necessary reaction condition. For methanol synthesis the CO bond remains intact. Hydrogen has two roles in catalytic syngas synthesis reactions. In addition to being a reactant needed for CO hydrogenation, it is usually used to reduce the metalized synthesis catalysts and activate the metal surface. Since the genesis of syngas conversion to fuels and chemicals a tremendous amount of research and development has been devoted to optimizing product yields and process

efficiencies. This includes the discovery of catalysts with optimized formulations containing the most active metals in combination with appropriate additives to improve activity and selectivity in a given process. Mechanistic studies have been conducted to interpret the fundamentals of specific conversion processes and measure the kinetic rates of key chemical reactions. Reactor design and engineering is another active research and development area of syngas conversion technology. Temperature control and stability in conversion reactors is a critical process parameter because of the large excess heat of reaction. Detailed process engineering and integration with respect to heat integration and syngas recycle to improve conversion efficiencies is used optimize commercial synthesis processes. Given the rich history of syngas conversion and the extensive research and development efforts devoted to this field of study, it is not surprising that a vast amount of literature is available that tracks the scientific and technological advancements in syngas chemistry. The purpose of this report is to review the many syngas to products processes and summarize the salient points regarding the technology status and description, chemistry, catalysts, reactors, gas cleanliness requirements, process and environmental performances, and economics. A diagram highlighting the specific processes reviewed in this report is shown in Figure 1.
Waxes Diesel Olefins Gasoline

isobutylene acidic ion exchange

Fe, Co, Ru




ThO2 or ZrO2

Syngas CO + H2
) os 3 ) 4 (Bu Ox CO ) 3P 3) 3 o( CO Ph HC Co( O)(P H h(C R

h ,R Co

ca CH rbon y 3O H latio Co +C n ,R O h, Ni

Mixed Alcohols


Acetic Acid

hom Co ologa tion


3 d l 2O pe /A do nO li3 /Z 3 ka r 2O Cu l 2O Al /C ; A O nO O/ Zn /Z Co Cu uO/ C oS 2 M

Di r U ect

zeolites MTO MTG

Olefins Gasoline

H2O WGS Purify

N2 over Fe/FeO (K2O, Al2O3, CaO)


the yn sis





Aldehydes Alcohols

M100 M85 DMFC

Figure 1: Diagram of Syngas Conversion Processes Covered in This Report

2.1 Introduction

The largest use of syngas is for hydrogen production (Wender, 1996). Hydrogen is produced as both a main product and as a by-product. Hydrogen producers often consume the product captively (e.g., ammonia producers, oil refineries, and methanol producers). The total amount of hydrogen consumed worldwide in 1999 was 15,864 billion ft3 and 20% of this was consumed in the U.S. (Suresh, et al, 2001). The majority of the worldwide hydrogen consumption (about 60%) was for ammonia production. This was followed by refining at 23% and methanol production at 9%, leaving only 8% as merchant hydrogen. Currently, hydrogen is primarily produced from methane. However, other hydrocarbon feedstocks include naphtha, heavy residues from petrochemical processes, coke oven gas, and coal. Electrolysis has also been used to produce hydrogen primarily where there is cheap electricity from hydropower stations. Presently, 77% of the worldwide hydrogen production comes from petrochemicals, 18 % from coal, 4 % from water electrolysis, and 1% from other sources (Hussinger, et al, 2000). The flexibility of hydrogen to be produced from any resource enables production from locally available, lowest-cost resources.


Technology Description

The dominant technology for hydrogen production today is steam methane reforming. The process can be divided into the following four steps: feed pretreatment, steam reforming, CO shift conversion, and hydrogen purification. For natural gas the only pretreatment required is desulfurization, which usually consists of a hydrogenator followed by a zinc oxide bed. After desulfurization, natural gas is fed to a reformer reactor, where it reacts with steam to produce CO, CO2, and H2. The reformer reactor is comprised of catalyst-filled tubes, surrounded by a firebox that provides the heat necessary for the endothermic reforming reaction, which operates at about 850C and between 1.5-3 MPa. The gas exiting the reformer is cooled to about 350C then undergoes the water gas shift (WGS) reaction (CO + H2O CO2 + H2) in a high temperature shift (HTS) converter. In the classical process, the gas is further cooled to about 220C and undergoes WGS in a low temperature shift (LTS) converter followed by CO2 scrubbing (e.g., monoethanolamine or hot potash). Finally, to remove trace amounts of CO and CO2 a methanation reactor is used (CO + 3H2 CH4 + H2O). The product hydrogen has a purity of 97-99%. Today, after the shift conversion step, hydrogen is purified using a pressure swing adsorption (PSA) unit to obtain purity greater than 99.99%. Since the PSA unit can easily remove CO and other components to produce a high purity hydrogen stream, most of todays hydrogen plants use only the HTS reactor (Leiby, 1994). However, the LTS reactor can increase the hydrogen yield slightly. The PSA off-gas, which contains unreacted CH4, CO, and CO2, plus unrecovered hydrogen, is used to fuel the reformer. This stream usually supplies 80-90% of the reformer heat duty supplemented by natural

gas to fulfill the remaining heat requirement (Leiby, 1994). Figure 2 is a block flow diagram of this process.

Figure 2: Block Flow Diagram of Hydrogen via Steam Methane Reforming


Hydrogen From Other Compounds

In addition to the direct production of hydrogen from gaseous hydrocarbon feedstocks, hydrogen can be produced from liquid energy carriers such as ethanol and methanol as well as from ammonia. Methanol reforming (CH3OH + H2O CO2 + 3H2) is practiced in Japan and to a small degree in Europe where there are no economical sources of syngas (Suresh, et al, 2001). Currently, there are about a dozen companies that are involved in building plants to produce hydrogen from methanol or ammonia.



Steam reforming of hydrocarbons involves the catalytic conversion of a hydrocarbon feedstock and steam to hydrogen and carbon oxides. Generally speaking, the chemical process for steam reforming of hydrocarbons is described by the following equation (Hussinger, et al, 2000): CnHm + nH2O (n+m/2)H2 + nCO Water gas shift is another important reaction that occurs in the reformer. For steam methane reforming the following two reactions occur in the reformer (Leiby, 1994): CH4 + H2O 3H2 + CO CO + H2O H2 + CO2 Hr = 49.3 kcal/mol Hr = -9.8 kcal/mol

Therefore, in this case 50% of the hydrogen comes from the steam. The reforming reaction is highly endothermic and is favored by high temperatures and low pressures. Higher pressures tend to lower the methane conversion. In industrial reformers the reforming and shift reactions result in a product composition that closely approaches equilibrium. The following side reactions produce carbon in the steam reformer:

2CO C(s) + CO2 CO + H2 C(s) + H2O CH4 C(s) + 2H2

(Boudouard coking) (CO reduction) (Methane cracking)

The reformer molar steam to carbon ratio is usually between 2-6 depending on the feedstock and process conditions. Excess steam is used to prevent coking in the reformer tubes. The shift reaction is exothermic and favors low temperatures. Since it does not approach completion in the reformer (usually there is 10-15 vol% CO, dry basis, in the reformer effluent), further conversion of CO is performed using shift conversion catalysts.



Conventional steam reforming catalysts are 10-33 wt% NiO on a mineral support (alumina, cement, or magnesia). Reforming catalyst suppliers include BASF, Dycat International, Haldor Topsoe, ICI Katalco and United Catalysts (Leiby, 1994). Heavy feedstocks tend to coke the reforming catalyst but promoters (potassium, lanthanum, ruthenium, and cerium) may be used to help this problem. These promoters increase steam gasification of solid carbon reducing coke formation but they also reduce the reforming activity of the catalyst (Leiby, 1994). For feedstocks heavier than naphtha, nickel-free catalysts containing mostly strontium, aluminum and calcium oxides have been successfully tested (Hussinger, et al, 2000). However, the methane content in the exiting gas is high requiring a secondary reformer. HTS catalyst has an iron oxide, chromium oxide basis while the major component in the LTS catalyst is copper oxide usually in a mixture with zinc oxide (Hussinger, et al, 2000). The HTS reactor operates in the temperature range of 300-450 C while the LTS is in the range of 180-270 C. Often the LTS reactor operates near condensation conditions. This catalyst is sensitive to changes in operating conditions. Typical catalyst lifetimes for both HTS and LTS catalysts are 3-5 years. Sulfur tolerant or dirty shift catalysts have also been developed that can handle larger amounts of sulfur. ICI Katalco makes dirty shift conversion catalysts that consist of cobalt and molybdenum oxides. They operate over a temperature range of 230-500 C. The controlling factors are the ratio of steam to sulfur in the feed gas and the catalyst temperature. For the methanantion catalyst, nickel oxide is used often in combination with chromium oxide.



A wide variety of reformer designs exist and can be used in various process configurations. The main components of the reformer furnace include an air/fuel combustion system, a radiant heat transfer section, and a convection section. The radiant section supplies heat to the catalyst tubes by combustion of the air/fuel mixture and the convection section recovers heat by cooling down the flue gases. Reformer


furnaces are not very efficient and only about half of the heat in the radiant section is absorbed by the furnace tubes. Generally, the feed gas flows up through the catalyst tubes. However, the reformer furnace can be side-, terrace-, top-, or bottom-fired (Johansen, 1992). Top-fired reformers are generally suited for larger production units. For small units of < 24 tubes, side-fired units are more economical. The terrace-fired reformer was developed in the 1960s is a variation of the side-wall design. Bottom-fired reformers are the least common of the four designs (Leiby, 1994). The shift converters and methanation reactor are fixed bed reactors.


Gas Cleanliness Requirements

Sulfur compounds are the main poison of reforming catalysts. Even at a concentration of 0.1 ppm the catalyst can begin to deactivate. To maintain a 3-year catalyst lifetime the sulfur concentration in the reformer feed gas should be less than 0.5 ppm (Leiby, 1994). Uranium oxide and chromium oxide are used as a promoter in certain reforming catalysts resulting in a higher tolerance to sulfur poisoning (Hussinger, 2000). Natural gas usually contains only small amounts of sulfur compounds generally in the form of H2S. In most SMR plants, prior to the steam reformer, there is a hydrogenator where sulfur compounds are converted to H2S. A desulfurizer follows this where the H2S in the feed gas is absorbed on a ZnO bed. Any remaining organic sulfur compounds and carbonyl sulfide are partially cracked and absorbed on the zinc oxide bed (Hussinger, et al, 2000). If the sulfur concentration in the feed gas is greater than 1% then the sulfur must be removed by chemical or physical scrubbing. The LTS catalyst is also very sensitive to sulfur. Additionally, chloride is a LTS catalyst poison. HTS catalyst can tolerate sulfur concentrations up to several hundred ppm, although the activity will decline. Other HTS catalyst poisons include phosphorus, silicon and unsaturated hydrocarbons in the presence of NOx (Hussinger, 2000). Prior to the PSA unit, entrained liquids (water and condensed hydrocarbons) must be removed because they will permanently damage the adsorbent, which is a mixture of activated carbon and zeolites. Cooling the product and installing a knock out drum with a mist eliminator prior to the PSA unit is usually sufficient.


Status/Technology Developers

BASF was the first company to develop steam methane reforming in the early 1900s (Nirula, 1995). In the mid-1930s to 1950s steam reforming was developed for heavier hydrocarbon feedstocks. Thereafter, energy efficiency improvements were made in this mature technology by the 1980s. There are many licensors for the various components of the steam reforming process. Table 3 is a list of licensors for steam reforming technology and Table 4 is a list of current commercial suppliers of PSA units.


Table 3: Catalytic Steam Reforming Licensors Company Headquarters Location ABB Lummus Global INc. Bloomfield, NJ, U.S.A. Air Products and Chemicals Allentown, PA, U.S.A. BOC Gases Murray Hill, NJ, U.S.A. Brown & Root, Inc. Alhambra, CA, U.S.A. John Brown Engineers & Constructors Ltd. London, England Caloric Grfelfing, Germany Exxon Research & Engineering Co. Florham Park, NJ, U.S.A. Haldor Topsoe A/S Copenhagen, Denmark Howe-Baker Engineers Tyler, TX, U.S.A. International Fuel Cells South Windsor, CT, U.S.A. JGC Corporation Tokyo, Japan KTI Group BV Zoetermeer, Netherlands Koch Process Technology Incorporated Roswell, GA, U.S.A. Linde AG Munich, Germany Lurgi AG Frankfurt, Germany Mahler AGS Stuttgart, Germany Messer Griesheim GmbH Krefeld, Germany The MW Kellogg Company Houston, TX, U.S.A. PETROBRAS Petrleo Basileiro S/A Ro de Janeiro, Brazil Uhde GmbH Dortmund, Germany From : (Suresh, et al, 2001 and Leiby, 1994) Table 4: Commercial Suppliers of PSA Units Company Headquarters Location LAir Liquide SA Paris, France Air Products and Chemicals St. Louis, MO, U.S.A. Bergbau AG Herne, Germany The BOC Group PLC Windlesham, United Kingdom Caloric Analgenbau Munich, Germany Howmar United Kingdom Linde AG Munich, Germany Questor Industries, Inc. Vancouver, BC, Canada Toyo Engineering Tokyo, Japan UOP, LLC Des Plaines, IL, U.S.A. From: (Suresh, et al, 2001)


Process Performance

About 65-75% of the CO and steam in the feed stream to the HTS reactor are converted to additional hydrogen and CO2 (Leiby, 1994). When a LTS reactor is used to convert additional CO to hydrogen about 80-90% of the remaining CO is converted to hydrogen resulting in an improvement in the hydrogen yield of about 5% (Leiby, 1994). For the PSA unit, the minimum pressure ratio between the feed and purge gas is about 4:1 and the purge gas pressure is typically between 17-20 psi to obtain a high recovery of hydrogen. Hydrogen recovery is usually 85-90% at these conditions and drops to 6080% at high purge gas pressures (55-95 psi) (Leiby, 1994). The PSA efficiency is also affected by adsorption temperature. Fewer impurities are adsorbed at higher


temperatures because the equilibrium capacity of the molecular sieves decreases with increasing temperature. Additionally, nitrogen is weakly adsorbed onto the adsorbent bed in the PSA unit, reducing the hydrogen recovery rate for the same purity. The hydrogen recovery may be reduced by as much as 2.5% for a 10-ppm nitrogen concentration in the PSA feed stream.

2.10 Environmental Performance

Hydrogen itself is a clean burning fuel. However, depending upon the feedstock used, its production can generate a considerable amount of CO2. Additionally, steam reformers produce NOx from fuel combustion. In California, selective catalytic reduction units are used in combination with low NOx burners to meet the states strict air regulations (Baade, et al, 2001). Controlling the emissions from less hydrogen-rich feedstocks is more difficult as the feedstock goes from natural gas to heavy fuel oil, coke, or coal. These feedstocks also contain other impurities such as sulfur and heavy metals.

2.11 Cost
Merchant hydrogen prices vary considerably depending on the volume and form of delivery. For large volume users, pipeline is the most economical followed by bulk liquid hydrogen delivery. For gaseous hydrogen produced from natural gas at a large-scale, central production facility at a pressure of around 400 psi, the plant gate price is about $5-$8/GJ. In general, contract prices for delivered hydrogen are not publicly available and they vary a great deal depending on the type of delivery, quantity required, and delivery distance. SRI states that liquid hydrogen list prices have been around $45/GJ, but that the average transaction prices are considerably below this (Suresh, et al, 2001). A typical price range for large-volume, bulk liquid consumers is $18-$24/GJ (Suresh, et al, 2001). Generally, as the feedstock goes from natural gas to light hydrocarbons to heavy hydrocarbons and then to solid feedstocks, there is an increase in processing difficulty and capital costs. Table 5 is a comparison of detailed biomass to hydrogen studies. Table 6 gives some costs for hydrogen from natural gas, coal, and biomass from other sources containing little or no details. All of these numbers are plant gate costs.


Table 5: Comparison of Biomass Based Hydrogen Studies

Study Cost year Biomass feed rate (BD tonne/day) Biomass cost Electricity selling price Net power (MW)

Hamelinck (2001) (a), (b) 2001 1,920 $2/GJ $34/dry tonne $0.03/kWh ranges from no excess to about 85 MW (@ 91% production cap) 91 184 (@ 91% production cap) $8-11/GJ Very detailed - Gives costs of individual equipment, operating costs, and other economic parameters and assumptions

Larson (1992) (a) 1991 1,650 $2/GJ $34/dry tonne $0.05/kWh (purchased) no excess

Spath (2000) (a) 2000 about 300-1,500 (3 plant sizes) $2.7/GJ $46.2/dry tonne $0.05/kWh (purchased) no excess (excess steam is produced and sold) 23 114 $14-21/GJ (for 15% IRR; gasification only) Somewhat detailed - Gives total installed capital costs, and other economic parameters and assumptions

H2 produced (tonne/day) Price of H2 (HHV basis, plant gate) Level of detail

134 - 188 $7-12/GJ Somewhat detailed Gives capital and operating costs

Notes: (a) Examined direct & indirect gasifiers and atmospheric and pressurized gasifiers. (b) Several configurations examined (varied type of gas cleaning, and H2 separation) Table 6: Hydrogen Cost Data From other Sources Resource Natural gas IEA (1999) $5.6/GJ (nat gas @ $3/GJ) $7/GJ (w/CO2 seq) Coal $10/GJ (coal @ $2/GJ) $13/GJ (w/CO2 seq) ----$1012/GJ $15/GJ (1987$) IEA (2001) $5.47.5/GJ McKinley, et al (1990) $7/GJ (1987$) British Government Panel (1999) $5.7/GJ (nat gas @ $3/GJ) $7/GJ (w/CO2 seq) $10.3/GJ (coal @ $2/GJ) $13.2/GJ (w/CO2 seq) ----$9.8719.3/GJ $9-10/GJ Grgoire Padr (1999) $5.977.46/GJ Amick (2003) -----



$7/GJ (Biomass @ $15/ton - 1987$) 14



2.12 References
Amick, P. Geosits, R., Kramer, S. (2003). 225th ACS National Meeting. New Orleans, Louisiana. March 23-27. Fuel Chemistry Preprints. 48 (1) pg. 199-204. Baade, W.F.; Parekh, U.N.; Raman, V.S. (2001). "Hydrogen." Kirk-Othmer Encyclopedia of Chemical Technology, John Wiley & Sons, Inc. British Government Panel on Sustainable Development . July 1999. Grgoire Padr, C.E.; Putsche, V. (1999). Survey of the Economics of Hydrogen Technologies. 57 pp.; NREL Report No. TP-570-27079. Hussinger, P.; Lohmller, R.; Watson, A.M. (2000). "Hydrogen." Ullmann's Encyclopedia of Industrial. Wiley-VCH Verlag GmbH & Co.KGaA. IEA Greenhouse Gas R&D Programme (1999) Hydrogen Today and Tomorrow. IEA Bioenergy Task 33: Thermal Gasification of Biomass, Spring 2001, Task Meeting. Johansen, T.; Raghuraman, K.S.; Hackett, L.A. design. Hydrocarbon Processing. (1992). Trens in hydrogen plant

Leiby, S.M. (1994). Options for Refinery Hydrogen. SRI Report No. 212. Menlo Park, CA. McKinley, K.R., Browne, S.H., Neill, D.R., Seki, A., Takahashi, P.K. (1990). Hydrogen Fuel from Renewable Resouorces. Energy Sources. Vol 12. pp. 105-110. Nirula, S.C. (1995). Synthesis Gas. SRI Report No. 148A. Menlo Park, CA. Suresh, B.; Gubler, R.; Sasano, T. (2001). "Hydrogen." Chemical Economics Handbook Product Review, SRI International, Menlo Park, CA. Reprot number 743.5000. Spath, P.L.; Lane, J.M; Mann, M.K; Amos, W.A. (April 2000). Update of Hydrogen from Biomass - Determination of the Delivered Cost of Hydrogen. NREL Milestone report. Wender, I. (1996). "Reactions of synthesis gas." Fuel Processing Technology 48(3): 189-297.


3.1 Introduction

Ammonia is the second largest synthetic chemical product. In 1996, worldwide ammonia production totaled 143 million tonnes (Appl 2003). Most of the ammonia produced is used to make fertilizers in the form of urea or ammonium salts (nitrate, phosphate, and sulfate). A small fraction of ammonia is used for the manufacture of organic chemical feedstocks for the plastics industry; polyamides, caprolactam, and others, and for the production of explosives (hydrazine, nitriles, etc.). Ammonia is also converted to nitric acid and cyanides. Ammonia is manufactured from nitrogen, fixed from the atmosphere, and hydrogen, produced mainly by steam methane reforming, in the catalytic process developed in the early 1900s by Fritz Haber and Carl Bosch using a promoted iron catalyst discovered by Alwin Mittasch (Satterfield 1991). The onset of World War I brought concerns about diminished worldwide nitrate supplies caused by an increase in the production of explosives. This led to the first industrial plant constructed at Oppau in Germany and commissioned in 1913. About 50% of the hydrogen produced from syngas processes is used for ammonia production (Wender). The vast amount of literature on ammonia synthesis over the past 100 years details process improvements, reactor engineering, catalyst developments, and mechanistic interpretation of the conversion of N2 and H2 to NH3 over Fe catalysts at high temperature and high pressure. Several reviews on ammonia production were consulted in writing this technology description (Kirth-Othmer)(Appl 2003; Jennings and Ward 1989; Satterfield 1991) and individual references to previous work can be found therein. For example, the Ammonia article in Ullmanns Encyclopedia of Industrial Chemistry contains over 1000 cited references (Appl 2003). Numerous books and entire journals such as Nitrogen and Ammonia Plant Safety are can be found that are dedicated to all aspects of ammonia synthesis. Commercial ammonia synthesis catalysts are basically the same Fe promoted catalysts first developed by Mittasch. The search for catalysts with high activity and long lifetimes at lower synthesis temperatures and lower pressures continues. Recent developments in ammonia synthesis catalysts include Ru-based formulations that are showing promise for commercial applications. Improved catalysts and process optimization are part of the continuing efforts to improve the economics of ammonia production by reducing process energy consumption. Energy consumption in ammonia plants has decreased nearly an order of magnitude since the early plants; from ~100 GJ/tonne NH3 to ~30 GJ/tonne NH3 (Jennings and Ward 1989). Part of this is a result of improved efficiencies in syngas (hydrogen) production by switching fuels, from coal to natural gas. Improved reactor design and optimized heat integration in the ammonia plant also contributes significantly to reducing energy consumption. One objective is to optimize heat recovery and minimize the capital and operating costs of the ammonia synthesis loop.


Technology Description


The basic process steps for industrial ammonia production are synthesis gas production, gas conditioning, compression, and ammonia synthesis. The goal of the synthesis gas production and gas conditioning operations is to provide a pure H2 stream for input into the ammonia converter. This requires a number of unit operations including feedstock pretreatment, syngas generation, CO conversion, and gas purification. These individual operations are summarized in the generic block flow diagram in Figure 3.
Steam Air


Gas Gas Pretreatment Pretreatment

Hydrodesulfurization ZnO adsorption Saturation H2 Recycle

Primary Reforming

Secondary Reforming
Partial Oxidation

CO Shift Conversion
High and Low Temperature




CO2 Removal
MEA/DEA Rectisol (CH3OH) Selexol K2CO3

Synthesis Gas Recycle


Ammonia Synthesis



Figure 3: Generic Ammonia Synthesis Process Flow Diagram For natural gas plants, feedstock pretreatment is essentially desulfurization prior to steam methane reforming. Feedstock preparation and handling is more complicated for solid fuel based systems. Secondary reformers are also used to further decrease the hydrocarbon content in the syngas. Following the reforming section, the syngas is cleaned and conditioned. Syngas cleaning is more involved for solid fuel based processes. Conditioning the cleaned syngas is required to increase the yield and purity of the hydrogen that is produced. CO is converted to H2 via the water gas shift reaction in high and low temperature shift reactors. Water is removed by cooling the syngas and CO2 is removed in an absorber-stripper process. Residual CO and CO2 in the product H2 is reduced to lower levels by what amounts to the reverse steam methane reforming process. A fraction the H2 is consumed in a methanation reactor to convert residual CO and CO2 to CH4. Purified H2 is then compressed and fed to the ammonia synthesis loop. Another configuration would be to replace the following process steps: low temperature shift, CO2 removal, methanation and possibly the secondary reformer with pressure swing adsorption. It is also possible to replace the reforming section with an autothermal reformer. However, currently the majority of ammonia plants follow the process scheme shown in Figure 3. Ammonia synthesis is achieved in catalytic converters at pressures ranging from 15-35 MPa at a minimum temperature of 430-480C. A maximum temperature of ~550C results from the balance between increasing reaction kinetics and decreasing the equilibrium NH3 concentration as the temperature increases. Thermodynamics suggests


that low process temperatures are favorable, however, the kinetics of the reactions dictates that high temperatures are required. The most effective catalysts from a process perspective are those that have the highest rate of conversion at the lowest temperatures (Jennings and Ward 1989). With effective gas cleanup and conditioning, catalyst lifetimes of 5-8 years can be achieved in commercial practice and up to 14 years in certain cases. Ammonia formation increases as the process pressure is increased. The optimum H2/N2 ratio is near 2 at high space velocities and approaches 3 at low space velocities as equilibrium becomes more dominant. Space velocities in commercial ammonia synthesis vary from 12,000/hr at 15 MPa to 35,000/hr at 80MPa (Appl 2003). Per pass conversion on the order of 10-35% is typically achieved. Ammonia is recovered from the synthesis loop by cooling the synthesis gas at process pressures to condense the ammonia. The liquid ammonia is separated from the gas, which is recycled back through the converter. Earlier plant designs used air or water-cooling for ammonia recovery. Modern synthesis plants use refrigeration to condense out the ammonia. The ammonia recovery is not extremely efficient so the recycled gas typically contains 4% NH3 plus any inert gases (Ar, He, CH4, etc.) that may be in the process stream. Purging some of the gas in the recycle loop before it is recycled minimizes inert gas concentrations. This purged gas also contains some product NH3 that is recovered, unconverted N2 and H2, plus the inert gases. The purged H2 is recovered and recycled back to the converter or used as a fuel in the process. Some of the inert gases also end up dissolved in the condensed NH3 product. Flashing the liquid ammonia in a pressure letdown step releases any dissolved gases that are separated from the product. Clearly, the ammonia synthesis process consists of many complex unit operations apart from the actual synthesis loop. The way in which these process components are combined with respect to mass and energy flow has a major influence on efficiency and reliability. Apart from the feedstock, many of the differences between various commercial ammonia processes lie in the way in which the process elements are integrated.



Ammonia synthesis occurs via the following exothermic chemical reaction: N2 + 3H2 2 NH3 H773K = -109 kJ/mol; H298K = -46.22 kJ/mol

Since the reaction is exothermic, maximum conversion at equilibrium occurs at high pressure and low temperature. Early attempts at ammonia production were hampered by a lack of understanding of thermodynamic equilibrium and the temperatures and pressures chosen for ammonia synthesis were unfavorable for reaction. Ammonia synthesis is typically performed over promoted Fe-catalysts. Numerous mechanistic studies based on the measurement of surface adsorbed species have determined that the rate-limiting step is the dissociative chemisorption of N2 on the


catalyst surface to form N atoms. Surface concentration of N atoms is governed by the equilibrium with gas phase H2 and NH3. The following rate expression for ammonia synthesis was developed by Temkin and Pyzhev in 1940 and has been widely applied to ammonia converter design and development (Jennings and Ward 1989) and references therein):

PH 3 PNH 2 3 2 k2 r = k1 PN 2 PNH 2 PH 3 3 2

+ = 1 , > 0

where the first term is the rate of ammonia formation and the second term is the rate of ammonia decomposition. This rate expression is valid over a very wide pressure range with fugacities substituted for pressures at very high pressure. The value for ranges between 0.4-0.75 for commercial processes where ammonia is recycled back to the reactor. Clearly, this expression is not valid for zero ammonia concentrations that are often encountered in the laboratory. For process using only pure reactants without ammonia recycle, the following rate equation applies:
1 r = kPH PN 2 2

= 0 .5



Mittasch screened over 2000 catalyst compositions before developing a promoted Fe catalyst that is, in principle, still commercially used today. Catalysts usually contain multiple promoters to improve the effectiveness and stability of the catalysts for long periods of time. Ammonia synthesis catalysts are made by cooling, casting, crushing, and sieving a fused melt of magnetite ore (Fe3O4) and promoter precursors. Promoters include Al2O3, K2O, CaO, MgO, etc. One of the functions of the promoters is to minimize Fe sintering during reduction and reaction. K2O is a structural promoter. Typical ammonia synthesis catalysts contain: 2.5-4% Al2O3, 0.5-1.2% K2O, 2.0-3.5% CaO, 0-1.0% MgO, and 0.20.5% SiO2 (naturally occurring in the magnetite ore). Commercial catalyst formulations are listed in Table 7. Finely divide iron is pyrophoric so Fe-catalysts are reduced to metallic form in situ in ammonia converters. Porosity is developed as the catalysts are reduced. During reduction, oxygen is removed from the magnetite crystal lattice without shrinkage. Metallic Fe is formed with essentially the same porous structure as the magnetite precursor (Jennings and Ward 1989). The production and preservation of this highly porous structure during reduction of the ammonia synthesis catalysts leads to highly active catalysts. The additions of structural promoters like Al2O3, MgO, and Cr2O3 is to facilitate the formation of highly porous metallic iron. Surface areas of freshly reduced promoted iron ammonia synthesis catalysts can be as high as 15-20 m2/g. Catalyst activity is not directly associated with Fe surface area but is also related to the complex interactions of the promoters. Alkali metal promoters in ammonia synthesis


catalysts are needed to attain high activity. Potassium is the most cost effective alkali promoter. Cs and Rb are effective promoters but more costly than K. Li and Na are poor promoters and not used commercially. Potassium is thought to interact with the Fe crystallites by increasing the dissociative sticking probability of N2 on the Fe sites and increasing catalytic activity. Dissociative adsorption of H2 on the Fe sites occurs rapidly at relatively low temperatures. Equilibrium is, therefore, attained for H2 adsorption/desorption at ammonia synthesis temperatures. Other catalysts are being developed such as promoted Ru on high surface area graphite supports. These recently commercial catalysts offer the possibility of greatly reducing synthesis temperatures and pressures with improved activity at high ammonia concentrations. The Kellogg Advanced Ammonia Process is based on a Ru catalyst that is claimed to be 40% more active than Fe catalysts.
Table 7: Summary of Commercial Ammonia Synthesis Catalyst Compositions Catalyst Vendor Name Haldor KM1 Topsoe KM1R* 35-4 35-8* Katalco S6-10 (ICI) S610R* 74-1 74-1R* Sud AS-4 Chemie As-4R* Fe oxides# 91-95 89-93 93.9 92.1 93.9 92.1 95.8 94.7 92 78%Fe 10% FeO Composition (wt%) K2O Al2O3 CaO SiO2 Other Operating Ranges T (C) P (bar) 340-550 88-588 350-530 BASF catalyst 60 100-600

5-9 7-11
0.7 0.9 0.7 0.9 0.6 0.8 3.0 3.8 3.0 3.8 2.1 2.6 2.4 3.2 2.4 3.2 1.5 1.9

promoters promoters In pre-reduced

*pre-reduced catalysts # For unreduced catalysts, Fe is in the form of magnetite (Fe3O4). catalysts the Fe is present as metallic Fe and FeO.



The basic design of an ammonia synthesis reactor is a pressure vessel with sections for catalyst beds and heat exchangers. Ammonia converters are classified by flow type (radial, axial, or cross flow) and cooling method (quench or indirect) used. Axial flow reactors are essentially top-to-bottom flow reactors. The design is comparatively simple, however, a fairly large pressure drop develops across the catalyst bed. A radial flow configuration feeds gas into an annular region between the reactor wall and the outer surface of the catalyst bed. Gas flows through the bed and exits out a central collection tube. This design minimizes pressure drop across a shallow bed with a large surface area. Radial flow converters tend to be tall vessels with relatively small diameter. The cross-flow configuration has a similar principle in that gas is introduced along one side of the reactor and is collected radially across the reactor by a collector on the other side.


Removing the heat generated from the exothermic synthesis reaction to maintain control of the reaction temperature is another design challenge. Quench converters are designed to introduce cool reactant gas at various points along the length of the catalyst bed. Interbed heat exchangers can also be used to remove heat at specific intervals along the bed, effectively separating the bed into multiple synthesis zones, or continuously along the bed with cooling tubes. These indirectly cooled designs allow for efficient reaction heat recovery that can be used in other parts of the process. Reactant gas can be circulated through the heat exchangers to preheat the ammonia synthesis gas or water is used to produce steam. A number of ammonia converters are commercially available. Seventeen different designs for ammonia converters are described by Zardi (Zardi 1982). The article in Ullmanns also contains figures and descriptions of several different ammonia converters (Appl 2003). Many designs are based on quench converters that have multiple catalyst beds in series and cold gas is introduced between the beds for temperature control. Early designs based on axial flow were easier to build but not as efficient as the more complex radial flow converters that have more recently been designed. A small subset of ammonia converters are depicted in Figure 4 and described below.

Figure 4: Various Ammonia Converter Designs


3.5.1 Topsoe radial-flow converter

One design uses 2 radial beds with quench gas injection between them. A similar radial flow design uses an inter-bed heat exchanger in the first catalyst bed. Cold ammonia synthesis gas is introduced from the bottom of the converter through the second catalyst bed then through the heat exchanger in the first catalyst bed. The cold gas flow through the second bed also provides indirect heat exchange. An additional heat exchanger is located at the bottom of the reactor to cool the reacted gases.

3.5.2 Kellogg 4-bed axial flow quench converter

This reactor consists of 4 catalyst beds held on separate grids. Quench gas is introduced in the spaces between the beds. A heat exchanger is located at the top of the vessel.

3.5.3 Kellogg horizontal converter

This reactor is a cross-flow converter design where gas flows through the catalyst bed perpendicular to the axis of the vessel. It is available in both quench and indirectly cooled versions.

3.5.4 Tube Cooled Converter

There are many variations in this design. Figure 4 shows a generic reactor.

3.5.5 ICI Lozenge Quench Converter

This is an additional reactor which is not shown in Figure 4. An axial flow, continuous catalyst bed is divided by lozenge distributors for quench gas addition.


Gas Cleanliness Requirements

In a carefully controlled ammonia synthesis processes, 10-year catalyst lifetimes are not uncommon. Loss of activity, however, does still occur by common deactivation mechanisms. Deactivation by sintering is a slow, gradual process that leads to loss of active area by Fe crystallization and oxidation. The activity of ammonia synthesis catalysts can also be reduced by trace impurities in the feed gas stream or solid impurities introduced during catalyst manufacture. Careful manufacturing processes are required to avoid the solid catalyst poisons, although, some are unavoidable because they are present as impurities in the natural magnetite starting material. Extensive purification of H2 and N2 is required to minimize gas phase catalyst poisons. Gas phase impurities can be separated into temporary and irreversible catalyst poisons. Oxygenated compounds like CO, CO2, H2O, and O2 are temporary, reversible ammonia catalyst poisons at low process temperatures. Removal of these oxygenates from the feed gas stream followed by re-reduction of the catalyst with pure H2 and N2 will restore catalyst activity. A typical maximum specification allowed for all oxygen containing compounds in the ammonia synthesis reactor feed is 10 ppm (total oxygen compounds). As temperatures increase above 520C water vapor can become an irreversible poison as Fe crystallites form. Sulfur and chlorine are the most efficient irreversible gaseous catalyst poisons. The sulfur is removed from the process gas stream before the primary reformer by a


hydrodesulfurization unit containing Co-Mo catalysts followed by a fixed bed ZnO absorber. Syngas from coal gasification can contain high levels of H2S and COS that must be removed. Chlorine attacks the alkali promoters to form volatile alkali chlorides that can vaporize at process temperatures. Chlorine also attacks the active metallic Fe sites converting them to Fe-chlorides. Trace levels of arsenic and phosphorous are also strong catalyst poisons. Table 8 summarizes the various catalyst poisons. Table 8: Summary of Ammonia Synthesis Catalyst Poisons Poison H2O CO CO2 O2 Sulfur (H2S) Chlorine As, P, Sb Type Rev Rev Rev Rev Irrev irrev Irrev Maximum concentration (ppm) 200 200 100 100 0.1 0.1 ---


Status/Technology Developers

In the early days the entire ammonia industry was based on coal as the feedstock. Today coal or coke (including coke oven gas) is no longer of major economic importance. In 1990, for example, only 13.5 % of the world ammonia capacity was based on this raw material. Apart from a few plants operating in India and South Africa, today, the majority of coal-based ammonia plants are found in China. As cheap petroleum feedstocks became available the industry moved toward natural gas as the main feedstock. More than 20 commercial ammonia synthesis processes are described in the article on Ammonia in Ullmanns Encyclopedia of Industrial Chemistry (Appl 2003). Commercial vendors include: ICI ( Linde (Ilg and Kandziora 1997) - (, Kellogg, Brown, and Root - ( and Haldor Topsoe ( Ammonia Casale ( Krupp-Uhde - (


Environmental Performance

There are safety aspects in handling and storing liquid ammonia because it is toxic. The release of CO2 has been an environmental concern related to ammonia production. Often times the CO2 is captured and sold, otherwise it is vented to the atmosphere.



The cost of ammonia is highly dependent on the feedstock price. In general, the price of ammonia has fluctuated between $100-$250/tonne since 1986 (Appl 2003). It is interesting to note that in 1969 the majority of the ammonia was produced in North America and Europe (54% of total worldwide production) and now Asia produces the largest fraction of ammonia (38% of the total in 1996). Table 9 gives the cost of ammonia from a few sources for different feedstocks.
Table 9: Comparison of Ammonia Economics Study Cost year Feedstock Feedstock cost Ammonia (tonne/day) Price of ammonia ($/tonne) Level of detail Dietz et al 1978 1981 Brava cane $9.77/dry ton 91-363 $213-314 Some details 1980 Natural gas na 900 $215 1980 Resid na 900 $271 1980 Coal na 900 $314 Appl 2003 1996 Natural gas $2.8/GJ na $206 Vacuum resid $2.0/GJ na $248 Coal $1.8/GJ na $362 Eggeman 2001 1990 Natural gas $1.9/GJ 1,000 $161 Minimal details

Minimal details


3.10 References
Appl, M. (2003). "Ammonia." Ullmann's Encyclopedia of Industrial Chemistry Release 2003, 6th Edition, Wiley-VCH Verlag GmbH & Co.KGaA. Dietz, C.K., Lutz, I.H., Roman, A. (1978) Ammonia Production from Brava Cane. Energy From Biomass and Wastes Symposium, August 14-18, Washington, D.C., 575604. Eggeman, T. (2001). "Ammonia." Kirk-Othmer Encyclopedia of Chemical Technology, John Wiley & Sons, Inc. Ilg, J., and Kandziora, B. (1997). "Linde ammonia concept." Ammonia Plant Safety & Related Facilities, 37, 341-352. Jennings, J. R., and Ward, S. A. (1989). "Ammonia Synthesis." Catalyst Handbook, M. V. Twigg, ed., Wolfe Publishing, Ltd., London, England, 384-440. Lauriente, D.H. (2001). "Ammonia." Chemical Economics Handbook Marketing Research Report, SRI International, Menlo Park, CA. Report number 756.6000. Satterfield, C. N. (1991). Heterogeneous Catalysis in Industrial Practice, Krieger Publishing Company, Malabar, FL. Zardi, U. (1982). "Review These Developments in Ammonia and Methanol Reactors." Hydrocarbon Processing, 61(8), 129-133.


4.1 Introduction

Methanol synthesis actually began in the 1800s with the isolation of wood alcohol from the dry distillation (pyrolysis) of wood. Research and development efforts at the beginning of the 20th century involving the conversion of syngas to liquid fuels and chemicals led to the discovery of a methanol synthesis process concurrently with the development of the Fischer-Tropsch synthesis. In fact, methanol is a byproduct of Fischer-Tropsch synthesis when alkali metal promoted catalysts are used. Methanol synthesis is now a well-developed commercial catalytic process with high activity and very high selectivity (>99%). Due to economic reasons, methanol is almost exclusively produced via reforming of natural gas (90% of the worldwide methanol, (Davenport 2002). However, a variety of feedstocks other than natural gas can be used. The long time interest in methanol is due to its potential fuel and chemical uses. In particular, methanol can be used directly or blended with other petroleum products as a clean burning transportation fuel. Methanol is also an important chemical intermediate used to produce: formaldehyde, dimethyl ether (DME), methyl tert-butyl ether (MTBE), acetic acid, olefins, methyl amines, and methyl halides, to name a few.


Technology Description

Currently, the majority of methanol is synthesized from syngas that is produced via steam reforming of natural gas. It can also be reformed using autothermal reforming (ATR) or a combination of steam methane reforming (SMR) and ATR. Once the natural gas is reformed the resulting synthesis gas is fed to a reactor vessel in the presence of a catalyst to produce methanol and water vapor. This crude methanol, which usually contains up to 18% water, plus ethanol, higher alcohols, ketones, and ethers is fed to a distillation plant that consists of a unit that removes the volatiles and a unit that removes the water and higher alcohols. The unreacted syngas is recirculated back to the methanol converter resulting in an overall conversion efficiency of over 99%. A generic methanol synthesis process flow diagram is shown in Figure 5.

Steam, O2 Natural Gas Desulphurization syngas (CO/CO2/H2) Compressor Purge Gas Methanol Converter Cooling and Distillation Methanol Steam Reforming

syngas recycle loop


Figure 5: Simplified Methanol Synthesis Process Flow Diagram

As is the case with Fischer-Tropsch synthesis, one of the challenges associated with commercial methanol synthesis is removing the large excess heat of reaction. Methanol synthesis catalyst activity increases at higher temperatures but so does the chance for competing side reactions (Supp and Quinkler 1984). Byproducts of methanol formation are: CH4, DME, methyl formate, higher alcohols and acetone. Catalyst lifetimes are also reduced by continuous high temperature operation and typically process temperatures are maintained below 300C to minimize catalyst sintering. Overcoming the thermodynamic constraint is another challenge in commercial methanol synthesis. The maximum per pass conversion efficiency of syngas to methanol is limited to about 25% (Wender 1996). Higher conversion efficiencies per pass can be realized at lower temperatures where the methanol equilibrium is shifted towards products; however, catalyst activities generally decrease as the temperature is lowered. Removing the methanol as it is produced is another strategy used for overcoming the thermodynamic limitations and improving the per pas conversion process efficiencies. Methanol is either physically removed (condensed out or physisorbed onto a solid) or converted to another product like dimethyl ether, methyl formate, or acetic acid. Controlling and dissipating the heat of reaction and overcoming the equilibrium constraint to maximize the per pass conversion efficiency are the two main process features that are considered when designing a methanol synthesis reactor, commonly referred to as a methanol converter. Numerous methanol converter designs have been commercialized over the years and these can be roughly separated into two categories adiabatic or isothermal reactors. Adiabatic reactors often contain multiple catalyst beds separated by gas cooling devices, either direct heat exchange or injection of cooled, fresh or recycled, syngas. Axial temperature profiles often have a sawtooth pattern that is low at the point of heat removal and increases linearly between the heat exchange sections. The isothermal reactors are designed to continuously remove the heat of reaction so they operate essentially like a heat exchanger with an isothermal axial temperature profile. The following section describes some of the many methanol converter designs found in the literature.



One of the more widely used commercial isothermal methanol converters is the Lurgi Methanol Converter (Haid and Koss 2001; Supp and Quinkler 1984). It is a shell and tube design similar to their Fischer Tropsch (FT) reactor. The tubes contain a proprietary Lurgi methanol catalyst (Cu/ZnO/Cr2O3 + promoters) and are surrounded by boiling water for reaction heat removal. These units operate at 50-100 bar and 230265C. Varying the pressure of the boiling water controls the reactor temperature. Byproduct steam is produced at 40-50 bar and can be used to run the compressor or to provide heat for the distillation process. The ICI Low pressure Quench Converter (Pinto and Rogerson 1977; Rogerson 1971; Rogerson 1984) is the most widely used adiabatic methanol converter. It is operated at 50-100 bar and 270C. The Cu/ZnO/Al2O3 catalyst is contained in a single bed supported by an inert material. Adding cold fresh and recycled syngas quenches the synthesis reaction and controls the reaction temperature. The gas is injected at


appropriate depths within the reactor through spargers called lozenges. There are horizontal layers of these lozenges that run across the converter from side to side and each has an outer surface covered with wire mesh and a central pipe that delivers the cold gas. ICI has an improved version of this reactor known as an ARC converter. The main technical difference is that instead of a single continuous catalyst bed, the bed is separated by distribution plates to form multiple consecutive catalyst domains. Kellogg, Brown, and Root (now Halliburton) has developed an adiabatic methanol converter that has multiple fixed bed reactors arranged in series and separated by heat exchangers. All of the recycled syngas is fed directly into the first reactor stage. The reactors have a spherical geometry to reduce construction costs and they also use less catalyst compared to the ICI Quench Converter. The Haldor-Topsoe Collect, Mix, Distribute (CMD) converter operates on a similar principle. Vertical support beams separate catalyst beds. The gas inlet at the bottom of the reactor provides fresh syngas that flows radially up through the first catalyst bed. At the top of the reactor, this first pass through gas is mixed with quench gas and distributed evenly so that it flows radially down through the second catalyst bed (Dybkjaer 1981). The cited benefit of this design is an increase in per pass conversion. Toyo Engineering Corporation has designed another version of a multistage radial flow methanol converter (MRF-Z) that uses bayonet boiler tubes for intermediate cooling. The tubes divide the catalyst into concentric beds. The Tube Cooled Converter is a reactor design that is simple to operate. The syngas enters at the bottom of the reactor where a manifold distributes the gas through tubes that act as a heat exchanger prior to the gas entering the catalyst bed. The Linde isothermal reactor, known as the Variobar converter, has the unique design feature of using coiled helical tubes embedded in the catalyst bed for heat removal. Spacers separate the multi-layer coils and boiling water is circulated through the tubes. Mitsubishi Gas Chemical in collaboration with Mitsubishi Heavy Industry has developed an isothermal reactor known as the MGC/MHI Superconverter (Takase and Niwa 1985). This reactor design uses double-walled tubes that are filled with catalyst in the annular space between the inner and outer tubes. The feed syngas enters the inner tubes and is heated as it progresses through the tube. The gas then passes downward through the catalyst bed in the annular space. Heat is removed on both sides of the catalyst bed by the boiling water surrounding the tubes as well as by the feed gas introduced into the inner tube (Tijm et al. 2001). A high conversion rate (ca. 14 % methanol in the reactor outlet) is cited for this reactor (Fiedler et al. 2003). Additional methanol converter designs include technologies using three phase systems similar in principle to the slurry reactors used for Fischer Tropsch synthesis (FTS). These technologies are collectively known as Liquid Phase Methanol Synthesis. ChemSystems, Inc. and Air Products and Chemicals, Inc. with Department of Energy funding developed a liquid-entrained catalytic reactor for converting low H2/(CO + CO2) ratio syngas into methanol known as LPMEOH (Bhatt et al. 1999; DOE 1992). The ability to convert low stoichiometric ratio (CO rich) syngas lends to using syngas from coal or biomass gasification for methanol production. The three phase slurry reactor provides better temperature control by uniformly dissipating the heat of reaction into the high heat capacity liquid. The LPMEOH process uses a supported Cu/ZnO catalyst (20-45 wt%) dispersed in circulating mineral oil with reactor temperatures of 225-265C and a pressure of 50 bar. Brookhaven National Laboratory has also developed a liquid phase methanol synthesis process that can be operated at lower pressures (< 5 atm)

with up to 90% conversion without syngas recycling (Tijm et al. 2001). This work has been done in collaboration with Amoco and in 1998 a 50 ml mini-pilot plant was successfully operated. Two other methanol conversion processes are based on systems where the product methanol is continuously removed from the gas phase by selective adsorption on a solid or in a liquid. The Gas-Solid-Solid Trickle Flow Reactor (GSSTFR) utilizes an adsorbent such as SiO2/Al2O3 to trap the product methanol (Herman 1991; Pass et al. 1990). The solid adsorbent is collected in holding tanks and the methanol is desorbed. In the Reactor System with Interstage Product Removal (RSIPR), a liquid solvent is used to adsorb the product methanol (Herman 1991).



Catalytic methanol synthesis from syngas is a classic high-temperature, high-pressure exothermic equilibrium limited synthesis reaction. The chemistry of methanol synthesis is as follows (Rostrup-Nielsen 2000): CO + 2H2 CH3OH Hr = -90.64 kJ/mol CO2 + 3H2 CH3OH + H2O Hr = -49.67 kJ/mol CO + H2O CO2 + H2 Hr = -41.47 kJ/mol For methanol synthesis, a stoichometric ratio, defined as (H2 CO2)/(CO + CO2), of about 2 is preferred. This means that there will be just the stoichiometric amount of hydrogen needed for methanol synthesis. For kinetic reasons and in order to control byproducts, a value slightly above 2 is normally preferred (Dybkjaer and Christensen 2001). Although methanol is made from mixtures of H2 and CO, the reaction is about 100 times slower than when CO2 is present (Wender 1996). Until as recently as the 1990s the role of CO2 in methanol synthesis was not clear. The water gas shift (WGS) activity of Cu catalysts is so high that it was difficult to deconvolute the role of CO and CO2 in methanol synthesis. Isotopic labeling studies have unequivocally proved that CO2 is the source of C in methanol (Ladebeck 1993; Wender 1996). CO is involved in the reverse WGS reaction to make H2 and CO2. CO2 is also thought to keep the catalyst in an intermediate oxidation state Cu0/Cu+ preventing ZnO reduction followed by brass formation (Ladebeck 1993). The proposed mechanism for catalytic methanol synthesis is believed to proceed through a long-lived formate intermediate. CO2 is adsorbed on a partially oxidized metal surface as a carbonate and hydrogenated. This intermediate is then hydrogenated in the rate limiting step. The copper catalyst sites have high activity for splitting the first C-O bond in CO2 that helps maintain the oxidation state of the active copper sites. At high concentrations, however, CO2 actually reduces catalyst activity by inhibiting methanol synthesis. The feed gas composition for methanol synthesis is typically adjusted to contain 4-8% CO2 for maximum activity and selectivity. Even though Cu has WGS activity, excessive amounts of H2O also leads to active site blocking that is poor for activity but improves selectivity by reducing byproduct formation by 50% (Chinchen et al. 1990).




The first high temperature, high-pressure methanol synthesis catalysts were ZnO/Cr2O3 and were operated at 350C and 250-350 bar. Catalyst compositions contained 20-75 atom% Zn. These catalysts demonstrated high activity and selectivity for methanol synthesis and proved robust enough to resist sulfur poisoning inherent when generating syngas from coal gasification. Over the years, as gas purification technologies improved (i.e., removal of impurities such as sulfur, chlorine, and metals) interest in the easily poisoned Cu catalysts for methanol synthesis was renewed. In 1966, ICI introduced a new, more active Cu/ZnO/Al2O3 catalyst that began a new generation of methanol production by a low temperature (220-275C), low pressure process (50-100 bar). The last high temperature methanol synthesis plant closed in the mid-1980s (Fiedler et al. 2003) and presently low temperature low pressure processes based on Cu catalysts are used for all commercial production of methanol from syngas. The synthesis process has been optimized to point that modern methanol plants yield 1 kg of MeOH /liter of catalyst/hr with >99.5% selectivity for methanol. Commercial methanol synthesis catalysts have lifetimes on the order of 3-5 years under normal operating conditions. The Cu crystallites in methanol synthesis catalysts have been identified as the active catalytic sites although the actual state (oxide, metallic, etc.) of the active Cu site is still being debated. The most active catalysts all have high Cu content, optimum ~60% Cu that is limited by the need to have enough refractory oxide to prevent sintering of the Cu crystallites. The ZnO in the catalyst formulation creates a high Cu metal surface area; it is suitably refractory at methanol synthesis temperatures and hinders the agglomeration of Cu particles. ZnO also interacts with Al2O3 to form a spinel that provides a robust catalyst support. Acidic materials like alumina are known to catalyze methanol dehydration reactions to produce DME. By interacting with the Al2O3 support material, the ZnO effectively improves methanol selectivity by reducing the potential for DME formation. Catalysts are typically prepared by the co-precipitation of metal salts such as nitrates or sulfates with a variety of precipitating agents. It is important to avoid contaminating methanol catalysts with metals that have FT activity (Fe or Ni) during synthesis. Incorporation of alkali metals in catalyst formulations should also be avoided because they catalyze the increased production of higher alcohols. Table 10 shows catalyst formulations for the many commercial manufacturers.
Table 10: Commercial Methanol Synthesis Catalyst Formulations Manufacturer IFP ICI BASF Shell Sud Chemie Dupont United Catalysts Haldor Topsoe MK-121 Cu at% 45-70 20-35 38.5 71 65 50 62 > 55 Zn at% 15-35 15-50 48.8 24 22 19 21 21-25 Al at% 4-20 4-20 12.9 12 31 17 8-10 Other Zr 2-18 Mg Rare Earth oxide 5 Patent date 1987 1965 1978 1973 1987 none found none found none found


Additional catalyst formulations have been presented in the literature with the purpose of improving per pass methanol yields (Klier 1982). The addition of Cs to Cu/ZnO mixtures has shown improved methanol synthesis yields. This only holds true for the heavier alkali metals as the addition of K to methanol synthesis catalysts tends to enhance higher alcohol yields. Cu/ThO2 intermetallic catalysts have also been investigated for methanol synthesis (Klier 1982). These catalysts have demonstrated high activitity for forming methanol from CO2-free syngas. Thoria based methanol catalysts deactivate very rapidly in the presence of CO2. Cu/Zr catalysts have proven active for methanol synthesis in CO-free syngas at 5 atm and 160-300C (Herman 1991). Supported Pd catalysts have also demonstrated methanol synthesis activity in CO2-free syngas a 5 110 atm at 260-350C.


Gas Cleanliness Requirements

Through the years, methanol catalysts have been optimized for maximum selectivity, space time yields, and long service life, but they still deactivate (Wender 1996). Under normal commercial operating conditions, CuO/ZnO methanol catalysts have quite long lifetimes, up to 2-5 years. Increasing the process operating temperature to maintain activity and yield, compensates for progressive catalyst deactivation. Copper has a very low activity for breaking C-O bonds or forming C-C bonds. As a result in methanol synthesis, wax formation is not usually a major problem nor is formation of coke from hydrocarbons (Twigg and Spencer 2001). Cu catalysts also do not have a strong tendency to catalyze FT reactions, or processes involving carbonium ion chemistry. Therefore, the major causes of methanol catalyst deactivation are poisoning and Cu sintering that results in loss of Cu surface area. Copper catalysts are extremely sensitive to site-blocking poisons such as reduced sulfur. To retain the long-term activity of Cu catalysts, it has been found empirically that the gas-phase sulfur concentration needs to be kept below 1 ppm and preferably below 0.1 ppm. High partial pressures of water vapor also inhibit the reaction, presumably by competitive adsorption on active sites. Interestingly enough, the catalysts formulation itself tends to reduce the effects of sulfur poisoning. The ZnO in the catalyst is very effective in limiting poisoning of copper catalysts because it acts as a poison sink, diverting the poisons away from the active Cu sites and scavenging the sulfur out of the gas stream by forming ZnS and ZnSO4. Commercial Cu-ZnO catalysts can absorb 0.4 wt% sulfur and still maintain 70% of their activity (Kung 1992). This helps to extend the catalyst lifetime. As it turns out, COS is not a gas phase methanol catalyst poison. Studies have shown that COS does not cause any deactivation over a concentration range of 0.6-9 ppm. The Cu/ZnO catalysts have no activity for COS adsorption or dissociation and at methanol synthesis conditions COS does not react with H2 to form H2S (Kung 1992). However, COS is a much more effective poison in the liquid phase methanol synthesis process than in the gas phase.


If reactor temperatures are not properly controlled, the highly exothermic methanol synthesis reactions can rapidly cause catalyst deactivation by sintering of the Cu crystallites. Reactor temperatures are maintained below 300C to minimize sintering. The presence of Cl in syngas has also been correlated with a greatly enhanced rate of sintering of copper crystallites most likely due to the formation of volatile copper chloride compounds. The mechanism of Cl poisoning can be explained by adsorbed Cl atoms that block or modify active catalyst sites. The Cl interacts with the Cu by forming CuCl2 that has a low melting point and high surface mobility that accelerates catalysts sintering. The presence of trace amounts of chlorine also causes enhanced sulfur poisoning probably because Zn chlorides also form removing these sites as possible sinks for sulfur capture. The limits on HCl content to avoid catalyst poisoning are more severe than for H2S poisoning, on the order of 1 ppb (Twigg and Spencer 2001). Other gas phase poisons in syngas that need to be avoided are metal carbonyls, particularly Ni and Fe carbonyls. Trace amounts of metal carbonyl compounds could be present in feed streams, especially in laboratory experiments where stainless steel tubing is used. Volatile metal carbonyl compounds deposit on the catalyst surface and block active sites. The deposition of Fe and Ni carbonyls affects the selectivity of the methanol catalysts. Ni increases the production of CH4 and Fe increases the production of FT products. Metal carbonyl concentrations should be below 5 ppb (Wender 1996). The purity of catalyst formulations is also critical to insure optimum activity and selectivity of methanol synthesis catalysts. It is very important to eliminate Ni and Fe impurities in catalyst precursors during catalyst formulation. These impurities in methanol synthesis catalysts reduce the selectivity and increase the yield of CH4 and FT hydrocarbon products. Alkali metals should also be avoided because these reduce overall catalyst activity and increase the production of higher alcohols. Unwanted SiO2 in the catalyst promotes wax formation and unreacted (non-spinel) alumina enhances DME formation. Other poisons to be avoided are As and P. Maintaining the low levels of sulfur in the syngas used to produce methanol is typically not a problem. Most syngas used for methanol synthesis is produced from steam reforming of methane over Ni catalysts that are very intolerant of sulfur. Therefore, sulfur compounds in the feedstock will normally be reduced to below 0.1 ppm by desulfurization prior to the primary reformer in hydrogen, ammonia, and methanol plants. Poisoning by chlorine or other compounds is also negligible in normal plant operation, because these are removed by the earlier stages of generation of synthesis gas. A summary of the gas cleanliness requirements for gas phase and liquid phase methanol production are in Table 11.


Table 11: MeOH Gaseous Contaminant Constraints

Sulfur (not COS) Halides Fe and Ni

ppmv <0.5 (<0.1 preferred) 0.001 0.005

Reference Kung, 1992 Twigg and Spencer 2001 Kung, 1992

LPMeOH (Novem 2000) Sulfur (including OCS) Total halides Acetylene Total unsaturates NH3 HCN Fe and Ni

ppmv 0.1 0.01 5 300 10 0.01 0.01


Status/Technology Developers

The first commercial methanol synthesis route was introduced by BASF of Germany in 1923 at pressures of 250-350 atm (3,675-5,145 psi). Plants operated at high pressures until low pressure routes were developed and by the early 1980s the majority of the producers had switched from the high pressure process to a low pressure one. This was done because the low-pressure process is more efficient, has a lower capital cost, and is cheaper to operate. ICI and Lurgi offer the two most widely used low pressure processes (Davenport 2002). Low pressure technologies are also offered by Mitsubishi Gas Chemical Company, Inc. of Japan, jointly by Haldor-Topsoe of Denmark and Nihon Suiso Kogyo of Japan, and by Halliburton. The following table gives the process conditions for each supplier:
Table 12: Methanol Technology Suppliers Technology supplier ICI (Synetix) T (C) 210-290 P (atm) 50-100 Notes Currently licenses 4 types of reactors: ARC, Tubular Cooled, Isothermal Linde, and Toyo (see reactor section for details) Tubular, isothermal reactor Tubular, isothermal reactor To date, no commercial plants based on this process. Spherical reactor geometry

Lurgi 230-265 50-100 Mitsubishi 240 77-97 Haldor-Topsoe & 260 48-300 Nihon Halliburton not found not found (Kellogg, Brown & Root) References: (Davenport 2002; Dybkjaer and Christensen 2001; Fiedler et al. 2003; Lee 1990; Takase and Niwa 1985; Tijm et al. 2001)

The worlds largest producer of methanol is Methanex and they account for 17% of the total global capacity (Davenport 2002). The next largest producer is SABIC which accounts for 6.5% of the total global capacity (Davenport 2002). Overall, as of January 1, 2002 worldwide annual capacity of methanol was 12.8 billion gallons (38 million tonnes, 9 million GJ) (Davenport 2002). In 2001, the U.S. consumed 28% of the total world methanol while North America produced 20% of the world supply (Davenport 2002). There are 18 methanol production plants in the U.S. (Lidderdale 2001).


Recently, Lurgi won a contract to build a methanol plant for National Petrochemical's subsidiary Zagros Petrochemical in Bandar Assaluye, Iran. The capacity will be 1.8 million tonnes/yr making it the largest methanol unit in the world and is expected to begin operation in 2004 []. Liquid phase methanol synthesis (LPMeOH) was invented in late 1970s (Hamelinck and Faaij 2001) and developed in the 1980s (Bhatt et al. 1999). In the 80s, Air Products and Chemicals, Inc. tested this technology at a pilot-scale facility with an output of 8 tpd of methanol in LaPorte, TX. Later, a 260 TDP LPMeOH demonstration plant was built in Kingsport, TN and began operation in 1997 (Bhatt et al. 1999). It was a joint venture between Eastman Chemical Company and Air Products and Chemicals, Inc. (Wender 1996).


Process Performance

Converting syngas to methanol in commercial units is limited by thermodynamic considerations to about 25% per pass (Wender 1996). Lurgis multitubular reactor produces close to 1 kg of methanol per liter of catalyst per hour. The actual per pass conversion of methanol is only 4-7vol% (Wender 1996). Removal of methanol would increase per pass conversions. Another way to increase the conversion is in situ dehydration of methanol to DME. Generally, 100 tons of methanol is converted to nearly 44 tons of hydrocarbons and 56 tons of water (Maiden 1983; Wender 1996).



Globally, formaldehyde production is the largest consumer of methanol (35% of worldwide methanol) (Davenport 2002). This is followed by MTBE at 25% and the third largest consumer is acetic acid at 9% (Davenport 2002). The following sections discuss the uses of methanol.

4.9.1 Formaldehyde
Formaldehyde was first commercially produced in 1890 in Germany by the catalytic partial oxidation of methanol with air over an unsupported Cu catalyst at atmospheric pressure (Satterfield 1991). By the beginning of the 20th century Cu was replaced by Ag as the catalyst of choice because yields were better. Iron molybdate catalysts also began to be used in the 1940s (Satterfield 1991). Formaldehyde produced over Ag catalysts proceeds via methanol dehydrogenation and partial oxidation according to the following reactions: (Chen and Masel 1995; Ruf and Emig 1999): CH3OH H2CO + H2 (dehydrogenation) Hr = +84 kJ/mol CH3OH + O2 H2CO + H2O (partial oxidation) Hr = -159 kJ/mol For the Ag catalyst process, a 50% methanol mixture in air is passed through a thin, fixed bed at 600C and slightly above atmospheric pressure. Metallic Ag is not active for


methanol decomposition. The active Ag site requires chemisorbed oxygen, which provides a site for methanol adsorption. For the iron molybdate [Fe2(MoO4)3] process, a lean methanol-air mixture is used. This reaction is more exothermic than the Ag process so heat removal is required (Satterfield 1991). The catalyst temperature is about 340-380C. Excessive reactor temperatures cause volatilization of molybdenum oxide, which reduces the selectivity of the process. A high oxygen partial pressure is required to maintain catalyst activity and in the presence of excess methanol catalyst activity is lost. Formaldehyde is produced commercially from methanol by the following three industrial processes (Reuss et al. 2003): (1) Partial oxidation and dehydrogenation with air in the presence of silver crystals, steam, and excess methanol at a temperature of 680-720 C. Otherwise known as the BASF process. The methanol conversion for this process is 97-98%. (2) The same as process (1) except in the presence of crystalline silver or silver gauze at a temperature of 600-650 C then the product is distilled and the unreacted methanol is recycled. The primary methanol conversion for this process is 77-87%. (3) Oxidation with only excess air in the presence of a modified iron/molybdenum/vanadium oxide catalyst at a temperature of 250-400 C. Also know as the Formox process. The methanol conversion for this process is 98-99%. Formaldehyde is used to make resins with phenol, urea, or melamine for the manufacture of various construction board products. The demand for formaldehyde is driven by the construction industry.

4.9.2 MTBE (Methyl tertiary-Butyl Ether)

In general, greater than 95% of the MTBE that is produced is used in the gasoline pool. MTBE is also used in the petrochemical industry for the production of isobutene. Additionally, it can be used in a number of chemical reactions including the production of methacrolein, methycrylic acid, or isoprene. One other use is solvent dewaxing of hydrocarbon oils (Peters et al. 2003). MTBE is produced by reacting isobutene with methanol in the presence of an acidic catalyst according to the following slightly exothermic reaction: i-C4H8 + CH3OH (CH3)3COCH3 Hr = -37 kJ/mol

The reaction temperature and pressure are 30-100C and 7-14 atm (100-200 psig), respectively (Schdlich et al. 2003) so that the reaction occurs in the liquid phase. Catalysts used are solid acids, zeolites (H-ZSM-5), and macroporous sulfonic acid ionexchange resins such as Amberlyst-15 (Collignon et al. 1999). A molar excess of methanol is used to increase isobutene conversion and inhibit the dimerization and oligamerization of isobutene. MTBE yields approaching 90% can be achieved at optimum reaction conditions. In general, most commercially available processes are similar consisting of a reaction and a refining section. In addition to the Snamprogetti and Hls (now Oxeno) processes, processes have been developed by Arco, IFP, CDTECH (ABB Lummus

Crest and Chemical Research Licensing), DEA (formerly Deutsche Texaco), Shell (Netherlands), Phillips Petroleum, and Sumitomo (Peters et al. 2003). Currently, there are over 140 MTBE plants with a total installed capacity of ca. 20 million tonnes/yr. The largest demand for MTBE has been primarily in the U.S. with North America consuming roughly 65% of the worldwide production of MTBE. In the U.S. consumption increased substantially with the Clean Air Act Amendments of 1990 but recently MTBE has come under environmental attack because of leaky tanks and groundwater contamination. California planned to ban MTBE from gasoline by January 1, 2003 but the ban has been pushed back a year. Recently, Alberta Envirofuels closed its 800,000 tonne/yr MTBE plant located in Edmonton, Canada (2002). They plan to convert the plant into an iso-octane facility (a gasoline additive that is an alternative to MTBE) at a cost of $50 million. With the reduction and possibly elimination of MTBE in gasoline, a worldwide decline in the demand for MTBE of about 3-5% is predicted (Davenport 2002).

4.9.3 Acetic Acid

The third most abundant chemical synthesized from methanol is acetic acid (CH3OH + CO CH3COOH). Vinyl acetate, acetic anhydride and terephthalic acid are all manufactured from acetic acid. Latex emulsion resins for paints, adhesives, paper coatings, and textile finishing agents are made from vinyl acetate. Acetic anhydride is used in making cellulose acetate fibers and cellulosic plastics. In terms of world production of acetic acid, approximately one-half comes from methanol carbonylation and about one-third from acetaldehyde oxidation (Wagner 2002). (Cheung et al. 2003) states that there is little doubt that the technology of the future is methanol carbonylation. In 1999, the worldwide demand for acetic acid was 2.8 million tonnes (6.17 109 lb) (Wagner 2002). Acetic acid is produced by the carbonylation of methanol with CO in the liquid phase using Co, Rh, or Ni catalysts promoted by iodine. The synthesis of acetic acid by carbonylation of methanol is one of the most important industrial applications of homogeneous catalysis (Roper 1991). Two processes have been developed based on this homogeneous catalytic synthesis. One is the BASF process, which is the older of the two and uses a Co/iodine catalyst at process conditions of 250C and 500-700 atm (3,000-10,000 psig) with 90% selectivity to acetic acid (from methanol). The second is the Monsanto process, which uses a Rh/iodine catalyst at process conditions of 180C and 30-40 atm with over 99% selectivity. This liquid phase process is initiated by the reaction of methanol with HI to yield methyl iodide. The active catalyst in the process is the metal carbonyl [RhI2(CO)2]- into which methyl iodide is inserted in the rate limiting step. Acetic acid is formed by the hydrolysis of the eliminated acetyl iodide species CH3COI that also regenerates HI (King and Grate 1985). The second, less severe Monsanto process has displaced the BASF process and all new installed capacity since 1973 is based on the Monsanto process (SRI 1994). The catalytic system is very corrosive and requires expensive steels for materials of construction. Complete recovery of the expensive Rh catalyst (~10-3 M) and recycle of HI (~0.1 M) is paramount to maintain favorable process economics. The high cost of Rh has lead to the search for other, lower cost, metals that could be used as acetic acid catalysts with similar performance.


4.9.4 Methanol to Gasoline (MTG)

The methanol to gasoline (MTG) process developed by Mobil Oil Corporation involves the conversion of methanol to hydrocarbons over zeolite catalysts. The MTG process, although considered the first major new synthetic fuel development since FTS, was discovered by accident in the 1970s by two independent groups of Mobil scientists trying to convert methanol to ethylene oxide and attempting to methylate isobutene with methanol over a ZSM-5 zeolite catalyst (Keil 1999). The MTG process occurs in two steps. First, crude methanol (17% water) is superheated to 300C and partially dehydrated over an alumina catalyst at 27 atm to yield an equilibrium mixture of methanol, dimethyl ether, and water (75% of the methanol is converted). This effluent is then mixed with heated recycled syngas and introduced into a reactor containing ZSM-5 zeolite catalyst at 350-366C and 19-23 atm to produce hydrocarbons (44%) and water (56%) (Hancock 1985). The overall MTG process usually contains multiple gasoline conversion reactors in parallel because the zeolites have to be regenerated frequently to burn off the coke formed during the reaction. The reactors are then cycled so that individual reactors can be regenerated without stopping the process, usually every 2-6 weeks (Kam et al. 1984). The MTG reactions may be summarized as follows (Wender 1996): 2CH3OH CH3OCH3 + H2O CH3OCH3 C2-C5 olefins C2-C5 olefins paraffins, cycloparaffins, aromatics The selectivity to gasoline range hydrocarbons is greater than 85% with the remainder of the product being primarily LPG (Wender 1996). Nearly 40% of the gasoline produced from the MTG process is aromatic hydrocarbons with the following distribution: 4% benzene, 26% toluene, 2% ethylbenzene, 43% xylenes, 14% trimethyl substituted benzenes, plus 12% other aromatics (Wender 1996). The shape selectivity of the zeolite catalyst results in a relatively high durene (1,2,4,5-tetramethylbenzene) concentration, 35% of the gasoline produced (MacDougall 1991). Therefore, MTG gasoline is usually distilled and the heavy fraction is processed in the Heavy Gasoline Treating unit to reduce the durene concentration to below 2%. This results in a high quality gasoline with a high octane number. The MTG process may no longer be a good option for gasoline production because the 1990 Clean Air Act Amendments limits the amount of aromatics in reformulated gasoline in order to reduce air toxics. In California, the amount of total aromatics is limited to 25 vol% (Owen and Coley 1995). The first commercial MTG plant came onstream in 1985 in New Zealand (Mobils Motunui plant) producing both methanol and high octane gasoline from natural gas . The plant produced 14,500 BPD of gasoline then in 1997, the gasoline manufacturing was abandoned and the plant produced only methanol. No information could be found regarding the reason for abandoning the manufacture of gasoline, but it is most likely due to economic factors such as the cost of oil and the market price of methanol. Today, this plant along with a nearby methanol plant at Waitara produces 2.43 million tonnes per year of chemical grade methanol for export [].


A fluid bed MTG plant was jointly designed and operated near Cologne, Germany by Mobil Research and Development Corp., Union Rheinische Braunkohlen Kraftstoff AG and Uhde Gmb (Keil 1999). A demonstration plant (15.9 m3/day) operated from 1982 to 1985. Although, no commercial plants have been built, the fluid bed technology is ready for commercialization.

4.9.5 Methanol to Olefins (MTO) and Methanol to Gasoline and Diesel (MOGD)
Along with the MTG process, Mobil developed several other processes for converting methanol to hydrocarbons based on zeolite catalysts. Since light olefins are intermediates in the MTG process, it is possible to optimize the methanol to olefins (MTO) synthesis. Higher reaction temperatures (~500C), lower pressures, and lower catalyst (acidity) activity favors light olefin production (Keil 1999). The rate of olefin production could be modified so that 80% of the product consists of C2 to C5 olefins rich in propylene (32%) and butenes (20%) with an aromatic rich C5+ gasoline fraction (36%) (MacDougall 1991; Wender 1996). The process can also be modified for high ethylene and propylene yield (>60%). Mobil also developed a methanol to gasoline and diesel (MOGD) process. Oligimerization, disproportionation and aromatization of the olefins produced in the MTO synthesis are the basis for the MOGD process. In the MOGD process, the selectivity of gasoline and distillate from olefins is greater than 95% (Keil 1999). One source gives the gasoline product from MOGD to be: 3 wt% paraffins, 94 wt% olefins, 1 wt% napthenes, and 2 wt% aromatics (Tabak and Yurchak 1990). A large-scale test run was performed at a Mobil refinery in 1981. Neither the MTO nor the MOGD process is currently in commercial practice (Wender 1996). However, UOP and HYDRO of Norway, does license their own methanol to olefins process where the primary products are ethylene and propylene [ and (Keil 1999)]. They use a fluidized bed reactor at 400 450C and achieve roughly 80% carbon selectivity to olefins at nearly complete methanol conversion (Apanel and Netzer 2002). The operating parameters can be adjusted so that either more ethylene is produced (48 wt% ethylene, 31 wt% propylene, 9 wt% butenes and 1.5 wt% other olefins) or else more propylene (45 wt% propylene, 34 wt% ethylene, 12 wt% butenes and 0.75 wt% other olefins).

4.9.6 TIGAS (Topsoe Integrated Gasoline Synthesis)

The Topsoe Integrated Gasoline Synthesis (TIGAS) process was developed by Haldor Topsoe with the intent of minimizing capital and energy costs by integrating methanol synthesis with the MTG step into a single loop without isolation of methanol as an intermediate (Keil 1999; Topp-Jorgensen 1987; Topp-Jorgensen 1988). This process was developed with the idea that future plants would be constructed in remote areas for recovery of low cost natural gas. In Mobils MTG process different pressures are preferred for syngas production, methanol synthesis, and the fixed bed MTG step. These pressures are 15-20 atm (221-294 psi), 50-100 atm (735-1470 psi), and 15-25 atm (221-368 psi), respectively (Wender 1996). The TIGAS process involves modified catalysts and conditions so that the system pressure levels out and separate


compression steps are not required. In order to do this a mixture of methanol and DME is made prior to gasoline synthesis then there is only one recycle loop which goes from the gasoline synthesis step back to the MeOH/DME synthesis step. A 1 MTPD demonstration plant was built in Houston, Texas in 1984 and operated for 3 years (ToppJorgensen 1988). The gasoline yield for the TIGAS process, defined as the amount of gasoline produced divided by the amount of natural gas feed and fuel, was shown to be 56.5 wt% (Topp-Jorgensen 1988). The TIGAS process, however, yields a lower quality gasoline compared to MTG with a lower octane number because of a reduced selectivity to gasoline range aromatics (MacDougall 1991).

4.9.7 DME (Dimethyl Ether)

DME is industrially important as the starting material in the production of the methylating agent dimethyl sulfate and is also used as an aerosol propellant. DME has the potential to be used as a diesel or cooking fuel, a refrigerant, or a chemical feedstock (Gunda et al. 1995; Peng et al. 1999b; Shikada et al. 1999). Commercial production of DME originated as a byproduct of high-pressure methanol production. DME is formed in a two-step process where first methanol is synthesized then it is dehydrated over an acid catalyst such as -alumina at methanol synthesis conditions. The DME reaction scheme is as follows (Hansen and Joensen 1991; Peng et al. 1999a): CO + 2H2 CH3OH (methanol synthesis) 2CH3OH CH3OCH3 + H2O (methanol dehydration) H2O + CO H2 + CO2 (water gas shift) NET reaction: 3H2 + 3CO CH3OCH3 +CO2 H = -21.6kcal/mole H = -5.6 kcal/mole H = -9.8 kcal/mole H = -58.6 kcal/mole

Note that one product in each reaction is consumed by another reaction. Because of the synergy between these reactions, syngas conversion to DME gives higher conversions than syngas conversion to methanol. (Shikada et al. 1999) gives the per pass and total conversion for the synthesis of methanol, methanol/DME and DME to be:
Table 13: Conversions for Methanol, Methanol/DME, and DME Conversion Per pass (%) Total (%) MeOH 14 77 MeOH/DME 18 85 DME 50 95

The optimum H2:CO ratio for DME synthesis is lower than that for methanol synthesis and ideally should be around one (Peng et al. 1999a; Peng et al. 1999b; Shikada et al. 1999). Recent improvements to the DME synthesis process involve the development of bifunctional catalysts to produce DME in a single gas phase step (i.e., one reactor) (Ge et al. 1998; Peng et al. 1999b) or the use of a slurry reactor for liquid phase dimethyl ether (LPDME) synthesis (Brown et al. 1991; Sardesai and Lee 1998).

4.9.8 M100 and M85

Neat methanol (M100) has many desirable properties that make it an attractive alternative to petroleum gasoline. It has a high heat of vaporization and relatively low heating value (about half of gasoline) that leads to a lower flame temperature compared to gasoline. This results in lower CO, NOx, and hydrocarbons emissions, but higher

formaldehyde emissions. Consequently, neat methanol has only found use as a motor fuel in high-performance racing engines and airplanes that have been fully modified and adapted to operation on methanol (Fiedler et al. 2003). This includes replacing plastic components in the fuel system, modifying the carburetor or fuel injection system, and preheating the fuel mixture. Methanol is miscible in water and in low concentration methanol mixtures phase separation can occur. Methanol is also more corrosive than gasoline, however, it is less toxic and is not carcinogenic. Methanol was used as a gasoline blending agent prior to the mid-1980s when the EPA prohibited the use of methanol in unleaded gasoline without the use of a cosolvent alcohol (Davenport et al. 2002). This was done primarily because methanol exhibits poor solubility in gasoline and has phase separation problems. However, there has been some success with M85, a mixture of methanol and gasoline (85 vol%/15 vol%, respectively). Methanol is most commonly used in a mixture with 15% gasoline for two reasons: (1) a methanol flame is colorless, so gasoline is added to give the flame some color so that a fire can be detected in the event of a vehicle crash and (2) neat methanol can cause cold-start problems in the winter, or vapor lock in the summer []. Although automobiles have been manufactured that operate on M85 and some fleets continue to use M85, this fuel has failed to develop on a large-scale. Contributing factors include adverse consumer perceptions coupled with a lack of sufficient infrastructure even though only small changes are required at gasoline fueling station (e.g., linings and seals in tanks, pumps, and dispensers) in order to handle M85 (Davenport et al. 2002). M100 is also being targeted for use in direct methanol fuel cells (DMFC) for fuel cell vehicle applications. DMFCs allow the use of methanol as fuel without requiring a fuel processor to extract hydrogen from the methanol. Significant progress has been made on finding better electrolyte material to prevent methanol crossover through the membrane. Operating temperatures of direct methanol fuel cells has increased to nearly 100C. Improved anode catalysts have facilitated methanol oxidation and eliminated the need for an on-board hydrogen reformer. Several automakers have developed methanol powered fuel cell vehicles and currently, DaimlerChrysler and Honda both have two prototype direct methanol fuel cell vehicles. In fact, in 2002 one of DaimlerChryslers vehicles (NECAR 5) completed a 3,000 miles cross-country trip across the U.S.

4.10 Environmental Performance

Methanol is stable under normal storage conditions. It is not subject to hazardous polymerization reactions, but can react violently with strong oxidizing agents. The greatest hazard involved in handling methanol is the danger of fire or explosion (English et al. 1995). When used in spark engines, the higher the concentration of methanol in a methanol/gasoline blend the lower the amount of CO, HC, and NOx in the exhaust emissions (Letcher 1983). In fact, the addition of tertiary butyl alcohol (TBA) and methanol may reduce the CO emissions up to 40 % and the hydrocarbon emissions up to 20 % in comparison with conventional gasoline (Schdlich et al. 2003).


As far as emissions are concerned, the MTG product has nearly the same emissions as gasoline from oil unless the methanol for the process is formed from biomass (Vermillion et al. 2001).

4.11 Cost of Methanol

From January 1993 to October 2001, U.S. spot prices for methanol ranged from a low of about $0.20/gallon to a high of about $1.50/gallon (Davenport 2002). In general, the price tended to be between $0.30/gallon ($5/GJ, LHV) and $0.70/gallon ($12/GJ LHV). Price fluctuations are often a result of oversupply but the price is also highly dependent on the availability of imports and the price of natural gas. The current price of methanol listed in Chemical Marketing Reporter (2003) for January 28th is $0.40/gallon ($7/GJ, LHV) given as a contract price at the point of production. However, on January 6th the price was listed as $0.62/gallon ($10/GJ, LHV) (2003). In producing methanol from natural gas, the majority of the capital cost is in the reforming step and gas-conditioning step. A capital cost breakdown for the ICI process is given as: desulfurization (2%), reforming/gas cooling (32%), steam production (14%), compression (24%), methanol synthesis (22%), and distillation (6%) (Wender 1996). Using a slurry reactor instead of tubular one reduces the capital cost and decreases the compression cost (i.e. lower pressure drop across the reactor). In fact, LPMEOH investment costs are expected to be 5-23% less than the gas phase process of the same capacity (Hamelinck and Faaij 2001). (Vermillion et al. 2001), claims that if methanol is formed from biomass, the cost would be $0.30/gallon to $0.55/gallon higher than if it were produced from natural gas. Table 14 gives cost information from several coal based methanol studies and Table 15 gives the cost for several biomass-to-methanol studies.


Table 14: Comparison of Coal Based Methanol Economic Studies Study (ChemSystems 1990) 1987 unclear $1.48/GJ $38.6/tonne 886 plus 833 MWe $6.6/GJ $0.45/gallon electricity sold at $0.37/kWh Level of detail Some details Gives total capital investment and operating costs, and other economic assumptions Somewhat detailed puts previous studies on same cost basis (Salmon 1986) Synthetic Fuels Associates, Inc. (1983) max MeOH once through 1982 23,054 (as received) $46.47/tonne 14,439 1982 9,662 (as received) $50.99/tonne 2,231 (Wham and Forrester III 1980) 1979 18,331 (MAF) $0.95/GJ 5,067 $8.2/GJ $0.56/gallon (Michaels 1979) (Bailey 1979)

Cost year Coal feed rate (tonne/day) Coal cost Methanol produced (tonne/day) Price of methanol (HHV basis)

1984 22,000-30,000 (as received) $0.95/GJ and $1.42/GJ 5,000-10,000 $7.2-$8.8/GJ $0.49-$0.60/gal

1979 4,100 (as received) $16.5/tonne (mine mouth) 1,995 $5/GJ $0.34/gal 1980$ and $8/GJ $0.53/gal 1984$ Some details Gives total capital investment and operating costs, and other economic assumptions

1979 15,767 (as received) $22/tonne 6,395 $8.2/GJ $0.56/gallon

plus 870 MWe $6.16/GJ $5.87-$9.43/GJ $0.42/gallon $0.40-$0.64/gal (price was set and IRR electricity sold at calculated to be $0.42/kWh marginal at 7%) Some details Gives total capital investment and operating costs, and other economic assumptions

Somewhat detailed Gives total capital investment, operating costs, and other economic assumptions

Somewhat detailed Gives capital cost by plant sections, operating costs, and other economic assumptions


Table 15: Comparison of Biomass Based Methanol Studies Study Cost year Biomass feed rate (BD tonne/day) Biomass cost (Hamelinck and Faaij 2001) (a), (b) 2001 1,920 $2/GJ $38/dry tonne $0.03/kWh no excess to about 100 MW (@ 91% production cap) none 613-970 (@ 91% production cap) $9-$12/GJ $0.58-$0.83/gallon (Novem 2000)
(a), (c)

(Larson 1992) (a) 1991 1,650 $2/GJ $38/dry tonne $0.05/kWh (purchased) no excess none 705-1,004 $9.62-$15.4/GJ $0.66-$1.05/gallon

1999 1,358 $3/GJ $55/dry tonne $0.067/kWh about 150 co-fires nat gas in gas turbine 57-196 $10.6-$13.6/GJ $0.72-$0.93/gallon

(Wyman et al. 1993) (a) 1989 1,814 $2.45/GJ $46/dry tonne not given no excess none 790-1,110 $10.01-$19.60/GJ $0.68-$1.34/gallon (@ 12% discount rate)

(1990) 1987 1,814 - 9,841 $2.45/GJ $46/dry tonne ($0-2.4/GJ sensitivity) $0.041/kWh no excess none 717 4,577 $10-$18.9/GJ $0.68-$1.29/gal (for 20% capital recovery factor)

(Mudge et al. 1981) 1980 1,800 $1.2/GJ $22/dry tonne ($0-154/dry tonne sensitivity) $0.03/kWh (purchased) no excess none 900 $8.2-$10.3/GJ $0.56-$0.70/gallon (sensitivity: $5.9-$22.7/GJ $0.40-$1.55/gal) Somewhat detailed - Gives capital cost by section, operating costs, and other economic parameters

Electricity selling price Net power (MW) Other fuels MeOH produced (tonne/day) Price of MeOH (HHV basis)

Very detailed - Gives Some details Somewhat Somewhat Somewhat costs of individual Gives costs of detailed - Gives detailed - Gives detailed - Gives equipment, operating major plant capital and capital and capital cost by costs, and other sections, and operating costs operating costs section and economic assumptions used and cash flow operating costs parameters and to determine sheet assumptions installed costs Notes: (a) Examined direct & indirect gasifiers and atmospheric and pressurized gasifiers. (b) Several configurations examined (varied type of gas cleaning, reforming, and gas or LPMeOH with and without recycle) (c) Examined once through LPMeOH with combined heat and power.

Level of detail


4.11.1 Cost of Formaldehyde

A study was performed to examine the three industrial routes for producing formaldehyde. The cost of production for each process is given in Table 16 below (Reuss et al. 2003). The BASF process is the most economical, followed by the Formox process. The key feature in the BASF process design is a liquid recirculation system that produces excess steam while simultaneous savings cooling water. For comparison, the price of formaldehyde given in Chemical Marketing Reporter for January 28th 2003 is $463/tonne ($0.21/lb).
Table 16: Comparison of Formaldehyde Production Costs

Process Cost of production ($/tonne) Return of capital investment ($/tonne) Cost of production with ROI ($/tonne)

BASF 345 33 378

Incomplete conversion & MeOH recovery 364 43 407

Formox 339 48 387

4.11.2 Cost of MTBE

From 1980 to 2001, the average U.S. Gulf Coast price for MTBE fluctuated between $0.63/gallon (1998) and $1.27/gallon (1981) (Davenport et al. 2002). MTBE is produced in three different manners: by refinery/petrochemical plants using by-product isobutylene, by merchant plants which isomerize n- and iso-butane and dehydrogenate isobutane to isobutylene, and by tertiary butyl alcohol (TBA) plants which react byproduct TBA from propylene oxide production with methanol (Lidderdale 2001). The first is the cheapest way to produce MTBE at a cost of $6,000 - $10,000/BPD of capacity. The second is the most expensive at a cost of $20,000 - $28,000/BPD of capacity and only 2 plants in the U.S. use the third manner to produce MTBE.

4.11.3 Cost of Acetic Acid

Chemical Marketing Reporter gives the current cost (January 28th 2003) of acetic acid to be $992/tonne ($0.45/lb).

4.11.4 Cost of MTG

The capital cost of the New Zealand MTG plant (14,450 BPD) was estimated to be $767 million dollars (mid-1980) with a total project cost of $1,475 million dollars (1985) (Bem 1985; Maiden 1983). The production cost including return on investment was estimated to be $0.96/gallon and the New Zealand government planned to sell the gasoline for $1.31/gallon (Bem 1985). Initial studies performed prior to design and construction of the plant came up with a gasoline cost of $0.74/gallon (Bem 1985). Table 17 is a comparison of three other studies.


Table 17: Comparison of Coal Based MTG Economic Studies Study Cost year Coal feed rate (tonne/day) Coal cost Gasoline produced (BPD) Price of gasoline (Lee et al. 1980) 1977 24,762 (as received) $7.7/tonne 23,065 $0.67-$0.87/gallon $5.1-$6.6/GJ (HHV) $5.5-$7.2/GJ (LHV) Not very detailed but refers to a more detailed report (Schreiner 1978) which gives overall capital costs and other details (Wham and Forrester III 1980) 1979 18,431 (MAF) $0.95/GJ 21,921 $1.25/gallon $9.5/GJ (HHV) $10.3/GJ (LHV) Some details Gives total capital investment, breakdown of operating costs, and other economic assumptions (Edwards et al. 1979) 1978 18,118 (as received) $5.5/tonne 15,000 $0.90/gallon $6.8/GJ (HHV) $7.4/GJ (LHV) Somewhat detailed Gives capital cost by plant sections, operating costs, and other economic assumptions

Level of detail

4.11.5 Cost of MTO

One study gives the cost of methanol, olefins via MTO, and gasoline via MTG to be $0.45/gal ($0.68/lb), $0.16/lb, and $0.59/gal ($0.99/lb), respectively (Gradassi 1998) for a natural gas price of $0.47/GJ. Overall, there are not many economic details given in the report.

4.11.6 Cost of DME

At a natural gas cost of $0.6/MMBTU, DME can be produced for $0.67/gallon (Chaumette et al. 1999). In comparison, their estimate for methanol production at this natural gas price is $0.35/gallon.


4.12 References
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Michaels, H. J. (1979). "Methanol from coal using the Koppers-Totzek process." Coal Technol. (Houston), 2nd(3), 163-78. Mudge, L. K., Robertus, R. J., Mitchell, D. H., Sealock, L. J., Jr., and Weber, S. L. (1981). "Economics of methanol and SNG production from biomass via catalytic gasification." Energy Biomass Wastes, 5th, 687-722. Novem. (2000). "Technical and Economic Databases - Biomass Energy Conversion Systems for the Production of Gaseous and/or Liquid Energy Carriers." Report No. GAVE-9915. Owen, K., and Coley , T. (1995). Automotive fuels reference book, Society of Automotive Engineers, Warrendale, PA. Pass, G., Holzhauser, C., Akgerman, A., and Anthony, R. G. (1990). "Methanol synthesis in a trickle-bed reactor." AIChE Journal, 36(7), 1054-60. Peng, X. D., Toseland, B. A., and Tijm, P. J. A. (1999a). "Kinetic understanding of the chemical synergy under LPDME conditions-once-through applications." Chemical Engineering Science, 54(13-14), 2787-2792. Peng, X. D., Wang, A. W., Toseland, B. A., and Tijm, P. J. A. (1999b). "Single-step syngas-to-dimethyl ether processes for optimal productivity, minimal emissions, and natural gas-derived syngas." Industrial & Engineering Chemistry Research, 38(11), 4381-4388. Peters, U., Nierlich, F., Schulte-Krne, E., Sakuth, M., Deeb, R., Laugier, M., Suominen, M., and Kavanaugh, M. (2003). "Methyl tert-Butyl Ethyl." Ullmann's Encyclopedia of Industrial Chemistry Release 2003, 6th Edition, Wiley-VCH Verlag GmbH & Co.KGaA. Pinto, A., and Rogerson, P. L. (1977). "Optimizing the ICI low-pressure methanol process." Chem. Eng. (N. Y.), 84(14), 102-8. Reuss, G., Disteldorf, W., Gamer, A. O., and Hilt, A. (2003). "Formaldehyde." Ullmann's Encyclopedia of Industrial Chemistry Release 2003, 6th Edition, Wiley-VCH Verlag GmbH & Co.KGaA. Rogerson, P. L. (1971). "The ICI [Imperial Chemical Industries] low pressure methanol process." Ingenieursblad, 40(21), 657-63. Rogerson, P. L. (1984). "Chapter 2-2: The ICI Low-Pressure Methanol Process." Handbook of Synfuels Technology, R. A. Meyers, ed., McGraw-Hill Book Company, New Tork. Roper, M. (1991). "Chapter 9 - CO Activation by Homogeneous Catalysts." New Trends in CO Activation, L. Guczi, ed., Elsevier, Amsterdam, 381-424. Rostrup-Nielsen, J. R. (2000). "New aspects of syngas production and use." Catalysis Today, 63(2-4), 159-164.


Ruf, S., and Emig, G. (1999). "Homogeneous catalysis by evaporated solids." 146(1-2), 271-278. Salmon, R. (1986). "Economics of Methanol Production from Coal and Natural Gas." ORNL-6091, Oak Ridge National Laboratory. Sardesai, A., and Lee, S. (1998). "Liquid phase dimethyl ether (DME) synthesis: A review." Reviews in Process Chemistry and Engineering, 1(2), 141-178. Satterfield, C. N. (1991). Heterogeneous Catalysis in Industrial Practice, Krieger Publishing Company, Malabar, FL. Schdlich, K., Schug, P., Fabri, J., Dabelstein, W., and Reglitzky, A. (2003). "Octane Enhancers." Ullmann's Encyclopedia of Industrial Chemistry Release 2003, 6th Edition, Wiley-VCH Verlag GmbH & Co.KGaA. Schreiner, M. (1978). "Research Guidance Studies to Asses Gasoline from Coal by Methanol-To-Gasoline and Sasol-Type Fischer-Tropsch Technologies." DOE Contract No. EF-76-C-01-2447, Mobil Research and Development Corp, Princeton, NJ. Shikada, T., Ohno, Y., Ogawa, T., Ono, M., Mizuguchi, M., Tomura, K., and Fujimoto, K. (1999). "Synthesis of dimethyl ether from natural gas via synthesis gas." Kinetics and Catalysis (Translation of Kinetika i Kataliz), 40(3), 395-400. SRI. (1994). "Acetic Acid and Acetic Anhydride." PEP 37B, SRI International, Menol Park, CA. Supp, E., and Quinkler, R. F. (1984). "The Lurgi Low-Pressure Methanol Process." Handbook of synfuels technology., R. A. Meyers, ed., McGraw-Hill, 113 - 131. Tabak, S.A., Yurchak, S. (1990). "Conversion of Methanol over ZSM-5 to Fuels and Chemicals." Catalysis Today, 6 (3), 307-327. Takase, I., and Niwa, K. (1985). "Mitsubishi (MGC/MHI) methanol process." CEER, Chemical Economy & Engineering Review, 17(5), 24-30. Tijm, P. J. A., Waller, F. J., and Brown, D. M. (2001). "Methanol technology developments for the new millennium." Applied Catalysis A, 221, 275-282. Topp-Jorgensen, J. (1987). "The Topsoe Integrated Gasoline Synthesis." Petrole et Techniques, 333, 11-17. Topp-Jorgensen, J. (1988). "Topsoe Integrated Gasoline Synthesis - The TIGAS Process." Methane conversion : proceedings of a symposium on the production of fuels and chemicals from natural gas, Auckland, April 27-30, 1987, D. M. Bibby, ed., Elsevier, Amsterdam ; New York, 293-305. Twigg, M. V., and Spencer, M. S. (2001). "Deactivation of supported copper metal catalysts for hydrogenation reactions." Applied Catalysis, A: General, 212(1-2), 161-174.

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5.1 Introduction

The production of higher alcohols from syngas has been known since the beginning of the last century. There are several processes that can be used to make mixed alcohols from CO and H2; including isosynthesis, variants of FTS, oxosynthesis involving the hydroformylation of olefins, and homologation of methanol and lower molecular weight alcohols to make higher alcohols. With the development of various gas-to-liquids processes such as Fischer-Tropsch and methanol synthesis, it was recognized that higher alcohols were by-products of these processes when catalysts or conditions were not optimized. Modified FT or methanol synthesis catalysts can be promoted with alkali metals to shift the products towards higher alcohols. Higher alcohol synthesis (HAS) is also optimized at higher temperatures and lower space velocities compared to methanol synthesis and with H2/CO 1 instead of 2 or greater. While other syngas-to-liquids processes were being commercialized, the commercial success of HAS has been limited by poor selectivity and low product yields. Single pass yields of HAS are on the order of 10% syngas conversion to alcohols with methanol typically being the most abundant alcohol produced (Herman 2000; Wender 1996). Methanol can be recycled to produce more higher alcohols or removed and sold separately. Despite these shortcomings, in 1913 BASF patented a process to synthesize a mixture of alcohols, aldehydes, ketones, and other organic compounds from CO and H2 over an alkalized Co oxide catalyst at 10-20 MPa and 300-400C (Herman 1991). Fisher and Tropsch developed the Synthol process for alcohol production in 1923. They used an alkalized Fe catalyst to convert syngas to alcohols at >10 MPa and 400-450C. Between 1935 and 1945 commercial mixed alcohol synthesis was performed with alkalized ZnO/Cr2O3 catalysts. The demand for mixed alcohol production from syngas decreased after 1945 with the increasing availability of petroleum and the desire for neat alcohols for manufacturing chemicals (Forzatti et al. 1991). Much of this early work on HAS is detailed in the review by Natta, Colombo, and Pasquon (Natta et al. 1957). The oil embargo of the 1970s provided the incentive for renewed interest in the synthesis and utilization of higher alcohols as a transportation fuel to be blended with gasoline. Mixed alcohols are a more attractive gasoline blending stock for octane enhancement compared to methanol. The octane rating is a measure of the fuels resistance to uncontrolled ignition by spurious sources such as hot spots in the combustion chamber (Mills and Ecklund 1989). The higher the octane number the less likelihood that a spurious ignition will occur. The undesirable properties of using neat methanol as a gasoline additive include high volatility, phase separation tendency when water is present, and incompatibility with certain engine fuel system components. Using mixed alcohols, containing methanol and higher alcohols, avoids these problems. Mixed alcohols have lower vapor pressure, better solubility with hydrocarbon components, improved water tolerance, and higher overall heating value compared to methanol. Also, when used as a diesel substitute at levels of 20-30 wt% higher alcohols, the calorific value, lubrication properties, and ignition properties are improved compared to pure methanol (Hhlein et al. 1991). Another important aspect is that it may be possible to use the existing fuel infrastructure.


The environmental impact of oxygenates in gasoline is another driver which increased the interest in HAS. The 1990 Clean Air Act mandated the seasonal use of oxygenated compounds in gasoline in specific regions of the U.S. to improve air quality by reducing volatile organic hydrocarbon emissions that lead to increased ground level ozone concentrations. Until recently, the oxygenate of choice as a gasoline additive was methyl tert-butyl ether (MTBE). The rapid increase in demand for MTBE led to concerns about shortages of petroleum-derived isobutene (Verkerk et al. 1999). Commercial synthesis of MTBE involves the etherification of isobutene with methanol over an acidic ion exchange catalyst. Recognizing an opportunity, research efforts increased to optimize the HAS process and catalysts to yield mixtures of methanol and isobutanol that could be combined in a dehydrative coupling reaction to produce MTBE from non-petroleum sources (Burcham et al. 1998; Verkerk et al. 1999). The optimum molar ratio of methanol to isobutanol for MTBE production is one. The isobutanol is dehydrated to yield isobutene that is then combined with methanol in the commercial MTBE synthesis process. These activities may not be important in the future considering the current environmental concerns being encountered with MTBE leaking into watersheds.


Technology Description

Currently there are no commercial plants that produce mixed alcohols in the C2 to C6 range. From a commercial viewpoint, the selectivity of HAS catalysts has not advanced to the point of commercial feasibility. However, the main process steps include synthesis gas production followed by gas clean up and conditioning, alcohol synthesis, and product purification.



The mechanism for HAS involves a complex set of numerous reactions with multiple pathways leading to a variety of products that are impacted by kinetic and thermodynamic constraints. No kinetic analysis of HAS has been published that is capable of globally predicting product compositions over ranges of operating conditions (Beretta et al. 1996). Depending on the process conditions and catalysts used, the most abundant products are typically methanol and CO2. The first step in HAS is the formation of a C-C bond by CO insertion into CH3OH. Linear alcohols are produced in a stepwise fashion involving the synthesis of methanol followed by its successive homologation to ethanol, propanol, butanol, etc. (Quarderer 1986). Therefore, the HAS catalyst should have methanol synthesis activity because methanol can be considered a recurrent C1 reactant. Branched higher alcohols are typically formed from modified methanol synthesis and modified FTS catalysts and straight chain alcohols are formed when alkalized MoS2 catalysts are used. The mechanism for HAS over modified high temperature methanol synthesis catalysts has been described as a unique carbon-chain growth mechanism that is referred to as oxygen retention reversal (ORR) aldol condensation with -carbon (adjacent to the alcohol oxygen) addition (Herman 2000). Individual reactions in HAS can be grouped into several distinct reaction types (Nunan et al. 1989):


1. Linear chain growth by C1 addition at the end of the chain to yield primary linear alcohols. 2. Beta addition between the C1 and Cn (n 2) to yield, for example, 1-propanol and branched primary alcohols such as 2-methyl-1-propanol (isobutanol) for n = 2 3. Beta addition between Cm (m = 2 or 3) and Cn (n 2). 4. Methyl ester formation via carboxylic acids formed from synthesized alcohols 5. Carbonylation of methanol to yield methyl formate Linear alcohols can proceed along the reaction path but branched alcohols are terminal products of the aldol condensation pathways because they lack the 2 -hydrogens required for chain growth (Hilmen et al. 1998). The general HAS reaction mechanism has the following overall stoichiometry (Hutchings et al. 1988; Wong et al. 1986): nCO + 2nH2 CnH2n+1OH + (n-1)H2O (Hr = -61.2 kcal/mol) with n typically ranging from 1 to 8 (Forzatti et al. 1991). The reaction stoichiometry suggests that the optimum CO/H2 = 2, however, the simultaneous occurrence of the WGS reaction means that the optimum ratio is closer to 1. The major reactions in HAS are: methanol synthesis, FT reactions, higher alcohol synthesis reactions, and the watergas shift (WGS) reaction (Xiaoding et al. 1987). The following is a list of some of these more important reactions described above that are associated with HAS. CO + 2H2 CH3OH CO + H2O CO2 + H2 CH3OH + CO CH3CHO + H2O CH3OH + CO + 2H2 CH3CH2OH + H2O CnH2n-1OH + CO + 2H2 CH3(CH2)nOH + H2O 2 CH3OH CH3CH2OH + H2O 2 CH3OH (CH3)2CO + H2O (CH3)2CO + H2 (CH3)2CHOH 2 CH3CHO CH3COOCH2CH3 Competing reactions: nCO + 2nH2 CnH2n + nH2O nCO + (2n+1)H2 CnH2n+2 + nH2O Methanol synthesis Water-Gas Shift CO beta addition - aldehydes Ethanol Homologation HAS Homologation Condensation/dehydration DME formation Branched iso-alcohols Methyl ester synthesis Olefins Paraffins

Methanol formation is favored at low temperatures and high pressures (Courty et al. 1990). At high pressures, HAS increases as the temperature is increased at the expense of methanol formation and minimizing hydrocarbon formation. To maximize higher alcohols, the H2/CO ratio should be close to the usage ratio, which is about 1. Lower H2/CO ratios favor CO insertion and C-C chain growth. In general, the reaction conditions for HAS are more severe than those for methanol production. To increase the yield of higher alcohols, methanol can be recycled for subsequent homologation provided the catalyst shows good hydrocarbonylation activity (Courty et al. 1990; Quarderer 1986). Unavoidably, the main reactions stated above produce H2O and CO2 as by-products. WGS plays a major role and depending on the catalysts shift activity, some chemical dehydration of alcohols can be undertaken in-situ to produce higher alcohols, esters, and ethers (Courty et al. 1990). Secondary reactions also produce hydrocarbons (HCs) including aldehydes and ketones (Courty et al. 1984; Courty et al.

1990). Also, frequently, substantial quantities of methane are formed (Roberts et al. 1992). Thermodynamic constraints limit the theoretical yield of HAS, and as in other syngas-to-liquids processes, one of the most important limitations to HAS is removing the considerable heat of reaction to maintain control of process temperatures (Courty et al. 1984). Compared to methanol, less alcohol product is made per mole of CO, more byproduct is made per mole of alcohol product, and the heat release is greater.



HAS catalysts are essentially bifunctional base-hydrogenation catalysts and are typically categorized into several groups based on their composition. Common to all HAS catalysts is the addition of alkali metals to the formulation. The activating character of alkali metal promoters is a function of their basicity. Alkali metals provide a basic site to catalyze the aldol condensation reaction by activating surface adsorbed CO and enhancing the formation of the formate intermediate. For the following, 4 groups of catalysts will be discussed (Herman 1991): 1) 2) 3) 4) Modified high pressure methanol synthesis catalysts alkali-doped ZnO/Cr2O3 Modified low pressure methanol synthesis catalysts alkali-doped Cu/ZnO and Cu/ZnO/Al2O3 Modified Fischer-Tropsch catalysts alkali-doped CuO/CoO/Al2O3 Alkali-doped sulfides, mainly MoS2

One of the major hurdles to overcome before HAS becomes an economic commercial process is improved catalysts that increase the productivity and selectivity to higher alcohols (Fierro 1993). To date modified methanol and modified FT catalysts have been more effective in the production of mixed alcohols; the sulfide-based catalysts tend to be less active than the oxide-based catalysts (Herman 2000). Rhodium (Rh) based catalysts are another group of catalysts that are not specifically used for HAS but have been developed for selective ethanol synthesis. Other C2 oxygenates (i.e., acetaldehyde and acetic acid) as well as increased levels of methane production are also synthesized over Rh-based catalysts (Nirula 1994). The high cost and limited availability of rhodium for ethanol synthesis catalysts will impact any commercialization of these synthetic processes for converting syngas to ethanol (Xiaoding et al. 1987).

5.4.1 Modified high catalysts





Alkali/ZnO/Cr2O3 - process conditions: 300-425C and 12.5-30 MPa - major product: Branched primary alcohols For non-alkalized catalysts containing Cu-Zn-Cr oxides, HAS yields were optimized for low Cr levels (15-21 wt% Cr). Chromia does not provide an active catalytic site but it is beneficial in small amounts because it acts as a structural promoter that increases the surface area of the catalyst and helps inhibit Cu sintering (Campos-Martin et al. 1995).


Methanol synthesis is fast compared to HAS over a 3% K2O/ZnCr catalyst, but is still equilibrium limited, even at high space velocities. CO2-rich feeds have been shown to inhibit HAS; a 3-fold decrease in C2+ alcohol production was measured for HAS at 400C with 6% CO2 (Tronconi et al. 1989). The productivity of primary alcohols was maximized over these alkali metal promoted high-temperature methanol synthesis catalysts with a CO2-free feed with H2/CO = 1 and CO conversion of ~ 5-20%. Nunan, et al. (Nunan et al. 1989) studied the effect of Cs-doping on Cu/ZnO catalysts. They determined that HAS product yields were maximized at 0.3-0.5 mol% Cs which translated to a surface concentration of 15-25%. The role of Cs-addition is to increase the ethanol synthesis rate that enhances the formation of higher alcohols. HAS conditions were 310C at 7.6 MPa with H2/CO = 0.45. Numerous papers by Hoflund and his group systematically describe the properties and effectiveness of promoted Zn/Cr HAS catalysts (Epling et al. 1997; Epling et al. 1998; Epling et al. 1999; Hoflund et al. 1997; Hoflund et al. 1999; Minahan et al. 1998a; Minahan et al. 1998b).

5.4.2 Modified low pressure/low temperature methanol synthesis catalysts

Alkali/Cu/ZnO(Al2O3) - process conditions: 275-310C and 5-10 MPa - major product: primary alcohols (Upper temperature of processes using copper catalysts is limited because of sintering) Many of the early commercial processes for HAS were based on alkali promoted, low temperature methanol synthesis catalysts. Lurgi developed the Octamix process in collaboration with Sd Chemie, who provided the catalyst for the process. The Octamix catalyst contained 25-40 wt% CuO, 10-18 wt% Al2O3, 30-45 wt% ZnO, and 1.7-2.5 wt% K2O with a Cu:Zn ratio of 0.4-1.9 with 3-18 wt% of a variety of oxidic promoters (Cr, Ce, La, Mn, or Th) (Xiaoding et al. 1987). With a starting gas composition of 25-30% CO, 08% N2, 0-5% CO2, 0-5% CH4, in a balance of H2; CO conversions were between 21-29% with 29-45% selectivity for C2+ alcohols and 17-25% CO2 selectivity at process conditions of 250-400C, 10 MPa, and a gas hourly space velocity of 1000-10,000/h (Xiaoding et al. 1987). Methanol is the most abundant oxygenated product (~ 80%) and the average carbon number of the oxygenated products is lower compared to products from modified high temperature methanol catalysts (Forzatti et al. 1991). The effectiveness of modified low temperature methanol synthesis catalysts has been detailed in the literature (Elliott and Pennella 1988; Nunan et al. 1989; Smith and Anderson 1984; Smith and Klier 1992).

5.4.3 Modified Fischer-Tropsch catalysts

Alkali/CuO/CoO-based catalysts - process conditions: 260-340C and 6-20 MPa - major product: Linear primary alcohols that follow ASF distribution The commercial HAS process sought and developed by the Institut Franais du Ptrole (IFP) and collaborators was based on modified Fischer-Tropsch catalysts. The IFP catalyst was a homogeneous mixed-oxide formulation containing Cu and Co on an alumina support as the active components for HAS modified with Zn and alkali metals. The catalysts were designed for HAS process conditions similar to low temperature methanol synthesis process conditions (5-15 MPa, T = 220-350C, H2/CO = 0.5-4 with CO2 also as a reactant) (Xiaoding et al. 1987). Patented IFP catalyst formulations have the following composition, on an element basis: 10-50% Cu; 5-25% Co; 5-30% Al; 1056

70% Zn; alkali/Al = 0-0.2; Zn/Al = 0.4-2.0; Co/Al = 0.2-0.75; Cu/Al = 0.1-3.0 (Xiaoding et al. 1987). The homogeneity of the catalyst correlates with good catalyst performance. The IFP process yields mainly saturated, straight-chained terminal alcohols that follow an Anderson-Shultz-Flory-type (ASF) distribution for chain growth. At optimal conditions, carbon conversion efficiency of CO and CO2 is between 5 and 30% and produces a liquid product containing 30-50% higher alcohols with hydrocarbons being the primary byproducts. The lack of long-term stability and low activity of these catalysts hinders the commercial application of this HAS process. Catalyst lifetimes have been quoted as long as 8000 h at the pilot-scale with little deactivation, caused mainly by coke formation and sintering that decreases the homogeneity of the catalyst (Xiaoding et al. 1987). Additional discussions of IFP-type catalyst studies have been presented in the literature (Courty et al. 1984; Courty et al. 1990; Courty et al. 1982; Dai et al. 1989; Dalmon et al. 1992).

5.4.4 Alkali-doped sulfides, mainly MoS2

Alkali/MoS2 - process conditions: 260-350C and 3-17.5 MPa - major product: linear alcohols In 1984, the Dow Chemical Company and Union Carbide Corporation independently disclosed a new catalyst system for converting syngas to linear alcohols. The catalysts were either supported or unsupported alkali-promoted MoS2 or Co/MoS2 (Herman 1991). Sulfide catalysts, such as MoS2, are well known hydrogenation catalysts, however, adding alkali metals as dopants shifts the products to alcohols rather than hydrocarbons. The role of the alkali metal is to suppress the hydrogenation activity of the Mo active site and to provide additional active sites for alcohol formation. Cesium is the most effective alkali promoter (Herman 2000), although K-addition has been extensively studied. The catalysts have 75-90% selectivity to higher alcohols from syngas with H2/CO = 1 with ~ 10% CO conversion efficiency (Herman 1991). Both higher alcohols and hydrocarbons formed over sufide catalysts follow a similar ASF molecular weight distribution. These catalysts have the unique property of being extremely resistant to sulfur poisoning, in fact, 50-100 ppm sulfur in the feed gas is required in the process to maintain the sulfidity of the catalyst (Courty et al. 1990). H2S in the feed gas is also thought to moderate the hydrogenation properties of the catalyst and improve selectivity of higher alcohols by reducing methanol production (Forzatti et al. 1991). Sulfide catalysts are also less sensitive to CO2 in the syngas but the presence of large amounts (>30%) of CO2 has been shown to retard the catalyst activity (Herman 1991). Moderate to low (7%) CO2 concentration in the syngas does not significantly impact CO conversion. The selectivity to higher alcohols versus methanol is reduced by the presence of even low levels of CO2 (Herman 1991) so CO2 removal is recommended for processes using sulfide catalysts. Adding Co to alkalized MoS2 catalysts increases the production of ethanol and other higher alcohols because Co promotes the homologation of methanol to ethanol (Forzatti et al. 1991). Activity of the sulfide catalysts can also depend on the catalyst support material (Avila et al. 1995; Bian et al. 1998; Iranmahboob and Hill 2002; Iranmahboob et al. 2002; Li et al. 1998).




Similar to other syngas conversion processes, one of the most important aspects of HAS is removing the large excess heat of reaction to maintain control of process temperatures, maximize yields, and minimize catalyst deactivation by sintering. HAS is performed in reactors that are similar to methanol and FT synthesis processes. Research and development is being conducted to investigate the use of slurry phase reactors for HAS. ChemSystems has conducted a pilot-scale study of isobutatnol synthesis in a slurry reactor using a Cs-promoted Cu/Zn/Al2O3 catalyst in hydrocarbon oil (40 wt% slurry) at 12.5 MPa and 350C (ChemSystems 1990; Herman 2000). Other HAS processes based on a double bed configuration have been explored. The idea is to optimize the production of methanol from syngas in the first reactor using a Cu-based catalyst at a lower temperature. The second reactor usually operates at a slightly higher temperature with a non-Cu Zn-chromite based catalyst to increase the yield of higher alcohols, particularly isobutanol by maximizing the C-C forming steps (Verkerk et al. 1999).


Gas Cleanliness Requirements

Gas cleanliness requirements for the modified FT and modified methanol synthesis catalysts are the same as those for the unmodified catalysts (see the gas cleanliness requirements section in the FT chapter and in the methanol chapter). Because the catalyst that Dow uses is already sulfided it is extremely resistant to poisoning by sulfur in the syngas (1984).


Status/Technology Developers

As stated previously, currently there are no commercial plants that produce mixed alcohols in the C2 to C6 range. Table 20 gives information regarding companies that have actively pursued mixed alcohols research. Generally, the H2/CO ratio is 1-1.4 except for the Snamprogetti which has a range of 0.5-3. In terms of commercialization, Dow, IFP and Snamprogetti are the most advanced in their process development (Nirula 1994).

5.7.1 Snamprogetti (also referred to as SEHT - Snamprogetti, Enichem, and Haldor Topsoe)
Snamprogetti and Haldor Topsoe developed this process jointly. A 12,000 tonne/yr pilot plant was started-up in 1982 in Pisticci, Italy (Olayan 1987). It was operated from 19821987 (Courty et al. 1990; Mills and Ecklund 1989; Nirula 1994). The process is similar to that for methanol production with a different catalyst and a higher operating temperature. In their proposed process, syngas is produced via POx of natural gas then mixed alcohols are synthesized in fixed bed adiabatic reactors. The crude alcohol mixture contains 20% water. The product is purified using 3 distillation columns. The first removes methanol and ethanol, the second is a stripping column to remove water, and the third recovers the C3+ alcohols using azeotropic distillation with cyclohexane (El Sawy 1990; Fox 1993). The final water content is below 0.1 wt% (Ricci et al. 1984).


The alcohol mixture was called MAS (Metanolo piu Alcoli Superiori methanol plus higher alcohols) (Mills and Ecklund 1989). It was blended into gasoline at about 5 vol% and marketed successfully as a premium gasoline known as SUPER E (El Sawy 1990; Mills and Ecklund 1989). Blends of 10 vol% were also tested in 13 different types of European cars and compared with conventional gasoline for driveability, acceleration, octane, exhaust emissions, and fuel economy (Ricci et al. 1984). For all of these categories except emissions, the alcohol/gasoline fuel blend was shown to be similar to conventional gasoline. The emissions varied depending on the type of car but in general there was a reduction in CO (as much as 40%) and unburned HC (roughly 18%), with slightly higher NOx emissions (Ricci et al. 1984). Extensive research was largely discontinued because of the availability of large amounts of cheap petroleum.

5.7.2 Dow
Dows process was first announced in 1984. They tested 2 high pressure reactors. One was a fluid bed and the other was a fixed bed (Nirula 1994). Dows catalyst has shown high productivity and a resistance to sulfur. The patented Dow Chemical catalysts contain free or combined Mo, W, or Re, a free or combined alkali or alkaline-earth metal, and optionally a support. A catalyst containing 21% Mo and 1.5% K on a carbon support demonstrated 23.4% CO conversion when a H2/CO mixture equal to 0.84 was passed over the catalyst at 262C and 7.2 MPa. The product distribution of the alcohols was: 25 mol% MeOH, 12.4 mol% EtOH, 6.8 mol% 1-propanol, 2 mol% 1-butanol, and 0.3 mol% C5 alcohols (Quarderer and Cochran 1984). However, in addition to the mixed alcohols, a substantial amount of CO2, methanol, and HCs have been produced (Xiaoding et al. 1987). The amount of methanol formed can be varied by adjusting the process conditions and the higher alcohols are predominantly straight chain C2-C5 alcohols. The water content is low (2-3 wt% given in (Quarderer 1986), and stated to be as low as 0.2% in1984 - Dow).

5.7.3 Lurgi (also referred to as Octamix)

While developing a low pressure methanol synthesis catalyst, a catalyst was discovered that had tendencies to yield higher alcohols. A 2 TPD demonstration plant was built in May 1990 at the Institute of Energy Process Engineering at the Research Centre Julich in cooperation with Lurgi (Goehna and Koenig 1989; Hhlein et al. 1991). This process differs from methanol synthesis only in the addition of CO2 removal and product purification. In Lurgis process scheme, syngas is produced via a combination of steam reforming and autothermal reforming then a tubular reactor is used for mixed alcohol synthesis. The water content of the crude mixture is only 1-2%. Thus a stabilizer column could be used instead of distillation or else molecular sieves to dry the product unless recovery of the methanol was desired (El Sawy 1990; Fox 1993). The final water content is low, 0.1-0.3 wt% (Goehna and Koenig 1989). The Octamix product was granted an EPA waver to be used as a gasoline additive on February 1, 1988 (El Sawy 1990).


Table 18: Companies and Information Regarding Mixed Alcohols Research Company Lurgi Union Carbide Sagami Research Center IFP Hoechst Snamprogetti Texaco (liquid phase system) Dow Country Germany U.S. Japan France Germany Italy U.S. U.S. Catalyst type Modified MeOH Rhodium Rhodium Modified MeOH Rhodium Modified MeOH Modified FT T (C) 250-420 300-350 200-300 260-320 275 260-420 220-240 P (psi) 725-1,450 1,000-2,500 735 850-1,450 1,455 2,610-3,822 6,615 H2/CO ratio 1-1.2 not found 1.4 1-2 not found 0.5-3 not found Products 53.5 wt% MeOH 41.9 wt% C2-C6 not found Mainly MeOH, EtOH, & CH4 30-50 wt% C2-C4 not found 20-40 wt%C2-C4 12-39 wt% non-alcohol oxygenates 30-70 wt% MeOH Conversion and Selectivity Information CO conversion = 20-60% CO selectivity to etoh = 60% CO conversion = 14% Selectivity to etoh up to 61% Selectivity to alcohols = 90% CO conversion = 12-18% Selectivity to alcohols = 70-75% CO selectivity to etoh = 74.5% CO conversion = 17% Selectivity to alcohols = 71% Syngas conversion = 40% Selectivity to products = 75%

Modified 299-310 1,500-2,000 1.1-1.2 CO conversion = 10-40% FT Selectivity to alcohols = 85% References: (Courty et al. 1990; Goehna and Koenig 1989; Herman 1991; Nirula 1994; Wender 1996).


5.7.4 IFP (also referred to as Substifuel)

IFPs process is in the laboratory stage and they built a 20 BPD unit in Chiba, Japan (Courty et al. 1990; El Sawy 1990). The IFP process scheme uses steam reforming followed by multibed quench synthesis reactors and distillation. The distillation section consists of 3 distillation columns: methanol distillation, extractive distillation with diethylene glycol (DEG), and distillation to recover the DEG (El Sawy 1990; Fox 1993). The final product contains about 0.2 wt% water.

5.7.5 EcaleneTM
Another mixed alcohol process, termed EcaleneTM, is currently at the bench scale but is being scaled up to a 500 gallon/day pilot plant. This is a collaborative project between Western Research Institute (WRI) and Power Energy Fuels, Inc. (PEFI), who developed the technology. Table 21 gives the composition of their alcohol mixture which can vary depending on the process parameters. The main alcohol constituent is ethanol. The synthesis reactor operates at 290-360C and 145-1,595 psi (Lucero et al. 2001).
Table 19: EcaleneTM Composition

Methanol Ethanol Propanol Butanol Pentanol Hexanol & higher

Current range (wt%) (Taylor 2002) 5-30 45-75 15 5 3 2

Target composition (wt%) (Lucero et al. 2001) 0 75 9 7 5 4


Environmental Performance

The evaporative emissions from fuel blends containing mixed alcohols (particularly methanol and ethanol) are typically much higher than emissions from straight gasoline. As a fuel additive, mixed alcohols have been shown to exhibit a reduction in CO, HCs, and nitric oxide emissions (Hhlein et al. 1991).



Table 22 summarizes the economics of higher alcohol synthesis from several studies which examine natural gas as the feedstock. Another source gives a breakdown of the capital cost to be 50% for syngas production, 29% for alcohol synthesis, 17% for CO2 removal, and 4% for alcohols fractionation (Courty et al. 1990). Economies of scale would improve the economics of mixed alcohol synthesis but it is also extremely important to reduce energy losses as well as overall heat and momentum transfer duty (Lange and Tijm 1996).

Table 20: Comparison of Mixed Higher Alcohol Economics From Natural Gas Study (ChemSystems 1990) 1986 $1.9/GJ and $0.47/GJ 50/50 550,000 $167.7/tonne $0.505/gallon $105.6/tonne $0.318/gallon (Salmon 1986) (Courty et al. 1984) (Ricci et al. 1984) (El Sawy 1990) (Bechtel 1998)

Cost year Natural gas cost ($/GJ) Methanol/higher alcohol mixture (wt%) Mixed alcohols produced (tonne/yr) Price of mixed alcohols

1989 $1.9/GJ and $0.47/GJ 65/35 for 70-90% selectivity 386,000 Production cost: ($1.9/GJ nat gas) $198-$255/tonne $0.60-$0.77/gallon ($0.47/GJ nat gas) $162-$199/tonne $0.49-$0.60/gallon

1983 $4.3/GJ (60-70)/(30-40) 600,000 Production cost: $313.9/tonne $0.95/gallon With 15% return: $400/tonne $1.21/gallon

1982 $0.8-$6.6/GJ 70/30 730,000 $110-$395/tonne $0.33-$1.19/gallon

1986 $2.8/GJ, $2.6/GJ and $0.47/GJ 70/30 and 50/50 635,000 730,000 Production cost: ($2.8/GJ nat gas) $282/tonne $0.774/gallon ($0.47/GJ nat gas) $205/tonne $0.562/gallon ($2.6/GJ nat gas) $0.90-$1.04/gallon

Not given $1/GJ 70/30 about 1,700,000 $225-$265/tonne $0.68-$0.80/gallon ($2.11/GJ nat gas) $279-$318/tonne $0.84-$0.96/gallon

Production cost: For $1/GJ nat gas $150/tonne $0.45/gallon Level of detail Few details Few details Few details Note: Production cost = no return on investment for these references.

Few details

Few details

Some details


5.10 References
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Dalmon, J. A., Chaumette, P., and Mirodatos, C. (1992). "Higher alcohols synthesis on cobalt based model catalysts." Catalysis Today, 15(1), 101-127. Elliott, D. J., and Pennella, F. (1988). "Mechanism of ethanol formation from synthesis gas over copper oxide/zinc oxide/alumina." Journal of Catalysis, 114(1), 90-9. El Sawy, A. H. (1990). "Evaluation of Mixed Alcohol Production Processes and Catalysts." NTIS. DE90010325. SAND89-7151, Mitre Corporation. Epling, W. S., Hoflund, G. B., Hart, W. M., and Minahan, D. M. (1997). "Reaction and surface characterization study of higher alcohol synthesis catalysts, II. Cspromoted commercial Zn/Cr spinel." Journal of Catalysis, 172(1), 13-23. Epling, W. S., Hoflund, G. B., and Minahan, D. M. (1998). "Reaction and surface characterization study of higher alcohol synthesis catalysts. VII. Cs- and Pdpromoted 1: 1 Zn/Cr spinel." Journal of Catalysis, 175(2), 175-184. Epling, W. S., Hoflund, G. B., and Minahan, D. M. (1999). "Higher alcohol synthesis reaction study. VI: effect of Cr replacement by Mn on the performance of Cs- and Cs, Pd-promoted Zn/Cr spinel catalysts." Applied Catalysis, A: General, 183(2), 335-343. Fierro, J. L. G. (1993). "Catalysis in C1 chemistry: future and prospect." Catalysis Letters, 22(1-2), 67-91. Forzatti, P., Tronconi, E., and Pasquon, I. (1991). "Higher Alcohol Synthesis." Catalysis Reviews-Science and Engineering, 33(1-2), 109-168. Fox, J. M., III. (1993). "The different catalytic routes for methane valorization: an assessment of processes for liquid fuels." Catalysis Reviews - Science and Engineering, 35(2), 169-212. Goehna, H., and Koenig, P. (1989). "The Octamix process." Indirect Liquefaction, Proc., Contract. Rev. Meet.(CONF-891131--DE90 008422), 59-83. Herman, R. G. (1991). "Chapter 7 - Classical and Non-Classical Routes for Alcohol Synthesis." New Trends in CO Activation, L. Guczi, ed., Elsevier, New York, 265349. Herman, R. G. (2000). "Advances in catalytic synthesis and utilization of higher alcohols." Catalysis Today, 55(3), 233-245. Hilmen, A. M., Xu, M., Gines, M. J. L., and Iglesia, E. (1998). "Synthesis of higher alcohols on copper catalysts supported on alkali-promoted basic oxides." Applied Catalysis A: General, 169(2), 355-372. Hoflund, G. B., Epling, W. S., and Minahan, D. M. (1997). "Higher alcohol synthesis reaction study using K-promoted ZnO catalysts. III." Catalysis Letters, 45(1,2), 135-138.


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6.1 Introduction

The synthesis of hydrocarbons from CO hydrogenation over transition metal catalysts was discovered in 1902 when Sabatier and Sanderens produced CH4 from H2 and CO mixtures passed over Ni, Fe, and Co catalysts. In 1923, Fischer and Tropsch reported the use of alkalized Fe catalysts to produce liquid hydrocarbons rich in oxygenated compounds in what was termed the Synthol process. Succeeding these initial discoveries, considerable effort went into developing catalysts for this process. A precipitated Co catalyst with 100 parts by weight Co, 5 parts by weight ThO2, 8 parts by weight MgO, and 200 parts by weight kieselguhr (silicious diatomatous earth) became known as the standard atmospheric pressure process catalyst. In 1936, Fischer and Pilcher developed the medium pressure (10-15 bar) FTS process. Following this development, alkalized Fe catalysts were implemented into the medium pressure FTS process. Collectively, the process of converting CO and H2 mixtures to liquid hydrocarbons over a transition metal catalyst has become known as the Fischer-Tropsch synthesis. Two main characteristics of Fisher-Tropsch synthesis (FTS) are the unavoidable production of a wide range of hydrocarbon products and the liberation of a large amount of heat from the highly exothermic synthesis reactions. Consequently, reactor design and process development has focused heavily on heat removal and temperature control. The focus of catalyst development is on improved catalyst lifetimes, activity, and selectivity. Single pass FTS always produces a wide range of olefins, paraffins, and oxygenated products such as alcohols, aldehydes, acids and ketones with water as a byproduct. Product distributions are influenced by temperature, feed gas composition (H2/CO), pressure, catalyst type, and catalyst composition. Product selectivity can also be improved using multiple step processes to upgrade the FTS products. In a review by Frohning, et al. (1954), it was cited that upwards of 4000 publications and a similar number of patents dealing with FTS could be found in the literature. Since then FTS has attracted an enormous amount of research and development effort. A comprehensive bibliography of FTS literature, including journal and conference articles, books, government reports and patents can be found in the Fischer-Tropsch Archive at This website is sponsored by Syntroleum Corporation in cooperation with Dr. Anthony Stranges, Professor of History at Texas A&M University and contains more than 7500 references and citations. This site has collected a bibliography of the large body of documents from the 1920's through the 1970s, which are important for researching the history and development of FTS and related processes as well as an up to date listing of the latest publications in this field. Many excellent reviews of FTS have been drawn upon for this report (Mills, 1993; Dry, 2002; and Dry, 1981) in an attempt to summarize, the chemistry (Frohning, et al, 1982), catalyst development (Bartholemew, 1991; Oukaci, et al, 1999; and Raje, et al, 1997), commercial processes (Dry, 2002, 1999, 1982; Senden, et al, 1992; Haid, et al, 2002, Van Nierop et al, 2000), reactor development (Adesini, 1986; Dry, 1988; and Dry and Hoogendoorm, 1981), and economics of FTS.



Technology Description

There are four main steps to producing FT products: syngas generation, gas purification, FT synthesis, and product upgrading. See Figure 7 for a generic process flow diagram. When using natural gas as the feedstock, many authors (Abbott and Crewdson, 2002; Rostrup-Nielsen, 2000; Vosloo, 2001; and Wilhelm, et al, 2001) have recommended autothermal reforming or autothermal reforming in combination with steam reforming as the best option for syngas generation. This is primarily due to the resulting H2/CO ratio and the fact that there is a more favorable economy of scale for air separation units than for tubular reactors (steam methane reforming - SMR). If the feedstock is coal, the syngas is produce via high temperature gasification in the presence of oxygen and steam. Depending on the types and quantities of FT products desired, either low (200 240C) or high temperature (300350C) synthesis is used with either an iron or cobalt catalyst. FTS temperatures are usually kept below 400C to minimize CH4 production. Generally, cobalt catalysts are only used at low temperatures because at higher temperatures a significant amount of methane is produced. Low temperatures yield high molecular mass linear waxes while high temperatures produce gasoline and low molecular weight olefins. If maximizing the gasoline product fraction, it is best to use an iron catalyst at a high temperature in a fixed fluid bed reactor. If maximizing the diesel product fraction, a slurry reactor with a cobalt catalyst is the best choice. The FT reactors are operated at pressures ranging from 10-40 bar (145580 psi). Upgrading usually means a combination of hydrotreating, hydrocracking, and hydroisomerization in addition to product separation.
Air, Oxygen, Steam Slurry (Co) or Tubular (Fe) Reactor

Waxes (> C20)


Coal, Natural Gas, or Biomass


Gas Cleanup and Conditioning

Clean syngas H2 and CO


Particulate Removal Wet Scrubbing Catalytic Tar Conversion Sulfur Scrubbing WGS etc.

CFB or FFB (Fe) Reactor

Olefins (C3 - C11)

Oligomerization Isomerization Hydrogenation

High T FTS


Figure 6: FTS General Process Flow Diagram




FTS has long been recognized as a polymerization reaction with the basic steps of 1) reactant (CO) adsorption on the catalyst surface, 2) chain initiation by CO dissociation followed by hydrogenation, 3) chain growth by insertion of additional CO molecules followed by hydrogenation, 4) chain termination, and 5) product desorption from the catalyst surface. Chemisorbed methyl species are formed by dissociation of absorbed CO molecules and stepwise addition of hydrogen atoms. These methyl species can further hydrogenate to form methane or act as initiators for chain growth. Chain growth occurs via sequential addition of CH2 groups while the growing alkyl chain remains chemisorbed to the metal surface at the terminal methylene group. Chain termination can occur at any time during the chain growth process to yield either an -olefin or an nparaffin once the product desorbs. The following is the FTS reaction: 1) CO + 2H2 --CH2-- + H2O Hr (227C) = -165kJ/mol (Haid, et al, 2000)

The water-gas shift (WGS) reaction is a secondary reaction that readily occurs when Fe catalysts are used. Combining reaction one (above) with reaction two (below) gives the net reaction for Fe catalyzed FTS (reaction 3). 2) 3) CO + H2O H2 + CO2 (Water-Gas Shift) 2CO + H2 --CH2-- + CO2 (net overall FTS)

The required H2 to CO ratio for the cobalt catalyst is 2.15 but since the iron catalyst performs WGS in addition to the FT reaction, the H2 to CO ratio can be slightly lower for the iron catalyst, 1.7 (Dry, 2002). Specific FTS products are synthesized according to the following reactions. 4) 5) 6) 7) CO + 3H2 CH4 + H2O (Methanation) nCO + (2n+1)H2 CnH2n+2 + nH2O (Paraffins) nCO + 2nH2 CnH2n + nH2O (Olefins) nCO + 2nH2 CnH2n+1OH + (n-1)H2O (Alcohols)

Another competing reaction that becomes important in FTS is the Boudouard reaction: 8) 2CO Cs + CO2

Carbon deposition on the catalyst surface causes catalyst deactivation. FTS is kinetically controlled and the intrinsic kinetics is stepwise chain growth, in effect the polymerization of CH2 groups on a catalyst surface. FTS product selectivities are determined by the ability of a catalyst to catalyze chain propagation versus chain termination reactions. The polymerization rates, and therefore kinetics, are independent of the products formed. The probability of chain growth and chain termination is independent of chain length. Therefore, selectivities of various hydrocarbons can be predicted based on simple statistical distributions calculated from chain growth


probability and carbon number. The chain polymerization kinetics model known as the Anderson-Shulz-Flory (ASF) model is represented by the following equation: Wn = n(1-)2n-1 Wn is the weight percent of a product containing n carbon atoms and is the chain growth probability. This equation is represented graphically in Figure 8 and clearly displays the predicted distributions for several products and product ranges of particular interest.
1.0 0.9 0.8 0.7


Weight Fraction

Wax (C35-C120) Gasoline (C5-C11) C2 C3 Diesel (C9-C25)

0.6 0.5 0.4 0.3 0.2 0.1 0.0 0.0











Probability of Chain Growth ()

Figure 7: Anderson-Shulz-Flory Distribution Irrespective of operating conditions, the FT reaction always produces a range of olefins, paraffins, and oxygenated compounds (alcohols, aldehydes, acids, and ketones). There are several unique aspects of FT products. Regardless of the product type, they are predominantly linear with high olefinicity. In fact, the paraffin-to-olefin ratio is lower than thermodynamically predicted. The olefins that do form are predominantly terminal (alpha). A considerable amount of monomethyl chain branches form and the degree of branching decreases as the chain length increases. Theoretically, only methane can be produced with 100% selectivity (Senden, et al, 1992). The only other product that can be produced with high selectivity is heavy paraffin wax. The gasoline product fraction has a maximum selectivity of 48% and the maximum diesel product fraction selectivity is closer to 40% and varies depending on the range of carbon numbers in the product cut. The variables that influence the distribution of these products are reactor temperature, pressure, feed gas composition, catalyst type and promoters.




Group VIII transition metal oxides are generally regarded as good CO hydrogenation catalysts. The earliest catalysts used for FTS were Fe and Co. Vannice has reported a relative activity of these metals for FTS (Adesina, 1996). In decreasing order of activity: Ru > Fe > Ni > Co > Rh > Pd > Pt. Ni is basically a methanation catalyst and does not have the broad selectivity of other FT catalysts. Ru has very high activity and quite high selectivity for producing high molecular weight products at low temperatures. Fe is also very active and has WGS activity. Fe readily forms carbides, nitrides, and carbonitrides with metallic character that also have FTS activity. Fe also has a stronger tendency than Ni or Co to produce carbon that deposits on the surface and deactivates the catalyst. Co tends to have a longer lifetime than Fe catalysts and does not have WGS activity, which leads to improved carbon conversion to products because CO2 is not formed. Co catalysts in FTS yield mainly straight chain hydrocarbons (no oxygenates like Fe). Although Ru is the most active FTS catalyst it is 3x105 times more expensive than Fe. Iron is by far the cheapest FTS catalyst of all of these metals. Co catalysts are 230 times more expensive than Fe but are still an alternative to Fe catalysts for FTS because they demonstrate activity at lower synthesis pressures so higher catalyst costs can be offset by lower operating costs. The 3 key properties of FT catalysts are lifetime, activity, and product selectivity. Optimizing these properties for desired commercial application has been the focus of FT catalyst research and development since the processes were first discovered. Each one of these properties can be affected by a variety of strategies including; use of promoters (chemical and structural), catalyst preparation and formulation, pretreatment and reduction, selective poisoning, and shape selectivity with zeolites. The performance of Co catalysts is not very sensitive to the addition of promoters. Early work demonstrated that the addition of ThO2 improved wax production at atmospheric pressure but had little effect at higher pressures. With Fe catalysts, however, promoters and supports are essential catalyst components. Since the discovery of FTS, potassium has been used as a promoter for Fe catalysts to effectively increase the basicity of the catalyst surface. The objective is to increase the adsorption of CO to the metal surface, which tends to withdraw electrons from the metal, by providing an electron donor. Potassium oxide addition to Fe catalysts also tends to decrease hydrogenation of adsorbed carbon species so chain growth is enhanced resulting in a higher molecular weight product distribution that is more olefinic. Potassium promotion also tends to increase WGS activity and lead to a faster rate of catalyst deactivation because of the increased rate of carbon deposition of the surface of the catalyst. Copper has also been successfully used as a promoter in Fe FTS catalysts. It increases the rate of FTS more effectively than potassium, however, decreases the rate of the WGS reaction. Copper has been shown to facilitate iron reduction. The average molecular weight of the products increases in the presence of copper but not as much as when potassium is used. Potassium promotion is not effective for Co catalysts. Catalyst preparation impacts the performance of Fe and Co catalysts. Fe catalysts can be prepared by precipitation onto catalyst supports such as SiO2 or Al2O3 or as fused iron where formulations are prepared in molten iron, cooled and crushed. The role of supports in Co catalysts is also important. Since Co is more expensive than Fe,


precipitating the ideal concentration of metal onto a support can help reduce catalyst costs while maximizing activity and durability. The combination of light transition metal oxides such as MnO with Fe increases the selectivity of light olefins in FTS. Fe/Mn/K catalysts have shown selectivity for C2-C4 olefins as high as 85-90%. Noble metal addition to Co catalysts increases FTS activity but not selectivity.



One of the challenges with FTS mentioned above is the removal of the large amount of excess heat generated by the exothermic synthesis reactions. Insufficient heat removal leads to localized overheating that results in high carbon deposition leading to catalyst deactivation. Methane formation also dominates at higher temperatures at the expense of desired FTS products. For large-scale commercial FTS reactors heat removal and temperature control are the most important design features to obtain optimum product selectivity and long catalyst lifetimes. Over the years, basically 4 FTS reactor designs have been used commercially (see Figure 9).
Syngas Intlet Gas Outlet

Gas Outlet

Gas Outlet

Hopper Catalyst Tubes


Slurry Bed
Steam Steam Standpipe

Wax Out

Catalyst Bed
Cooling Water Cooling Water Syngas Inlet

Gas Outlet

Syngas Inlet

Wax Outlet

Syngas Inlet

Multi-tubular (ARGE) Fixed Bed

Circulating (Synthol) Fluidized Bed

Fixed (Sasol Advanced Synthol) Fluidized Bed

Fixed Slurry Bed

Figure 8: Types of Fischer-Tropsch Synthesis Reactors One of the earliest FTS reactor designs was the fixed bed tubular reactor. After many years of development Ruhrchemie and Lurgi have refined this concept into what is known as the ARGE high capacity FT reactor that has been used since the inception of Sasol I. They contain 2,000 tubes filled with Fe catalyst immersed in boiling water for heat removal. The water bath temperature is maintained in the reactor by controlling the


pressure. Syngas is introduced into the top of the reactor, flows through the tubes, and the products exit at the bottom of the reactor. Wax accounts for 50% of the products. Conversion efficiencies are on the order of 70%. The reactor is operated at 20-30 bar at an operating temperature of 220-260C. Additional temperature control is obtained by using high gas velocities and gas recycling. Recycled gas typically is 2.5 times the amount of fresh syngas introduced into the reactor. Catalyst lifetimes are around 70-100 days (Wender, 1996) and catalysts removal can be quite difficult. High temperature circulating fluidized bed reactors have been developed for gasoline and light olefin production. These reactors are known as Synthol reactors and operate at 350C and 25 bar. The combined gas feed (fresh and recycled) enters at the bottom of the reactor and entrains catalyst that is flowing down the standpipe and through the slide valve. The high gas velocity carries the entrained catalyst into the reaction zone where heat is removed through heat exchangers. Product gases and catalyst are then transported into a large diameter catalyst hopper where the catalyst settles out and the product gases exit through a cyclone. These Synthol reactors have been successfully used for many years, however, they have a number of limitations. They are physically very complex reactors that involve circulation of large amounts of catalyst that leads to considerable erosion in particular regions of the reactor. Synthol reactors have also reached their design limit in terms of capacity and throughput (7,500 BPD capacity, (Lutz, 2001)). The fixed fluidized bed Sasol Advanced Synthol reactor has replaced the circulating fluidized bed Synthol reactor. Gas is introduced through a distributor and bubbles up through the catalyst bed. Heat is removed by an internal heat exchanger immersed in the catalyst bed. These new reactors are half the cost and size of the circulating reactors for the same capacity, have better thermal efficiency with a less severe temperature gradient and a lower pressure drop across the reactor. Operating costs are considerably lower and there is greater process flexibility (in terms of product distribution) and the possibility for scale-up to 20,000 BPD (Lutz, 2001). Process conditions in the fixed fluidized bed reactors are similar to those established in the Synthol reactors. Another reactor design is the low temperature slurry reactor, which has been considered since Kolbels pioneering work in the 1950s (Dry, 1996; and Dry, 2002). These 3-phase reactors consist of a solid catalyst suspended and dispersed in a high thermal capacity liquid (often the FT wax product). Syngas is bubbled through the liquid phase achieving excellent contact with catalyst while keeping the catalyst particles dispersed. Slurry reactors are optimized at low temperatures for FT wax production with low methane production. Compared to the fluidized bed reactors, the liquid slurry bed offers the advantages of better temperature control, lower catalyst loading, and significantly lower catalyst attrition rates. The improved isothermal conditions in slurry bed reactors allows for higher average reactor temperatures leading to higher conversions to products. Slurry bed reactors also cost 75% less than the much more complex multitubular fixed bed reactors. These reactors have only recently been put into commercial FT production primarily because one of the technical barriers, which required considerable development, was reliable catalyst separation from the FT waxes.



Gas Cleanliness Requirements

FTS catalysts can lose activity as a result of 1) conversion of the active metal site to an inactive oxide site, 2) sintering, 3) loss of active area by carbon deposition, and 4) chemical poisoning. Some of these mechanisms are unavoidable and others can be prevented or minimized by insuring that the impurity levels in the incident syngas are acceptable for the given process. Coke deposition is the most important mode of catalyst deactivation that can be impacted by addition of promoters to catalysts and reaction temperature and pressure. This mode of catalyst deactivation is largely unavoidable and FTS processes must be operated in a manner that the decreasing output from coke deposition is balanced with the economic considerations of catalyst regeneration or replacement. In general, because of its high activity, the coke deposition rate is higher for Fe catalyst than Co catalyst. Consequently, Co catalysts have longer lifetimes. One of the more controllable modes of catalyst deactivation is that induced by poisoning of the active sites by impurities in the syngas. By far the most abundant, important, and most studied FTS catalyst poison is sulfur. Sulfur is present in both natural gas and coal and during steam reforming or gasification gets converted primarily to H2S plus other organic sulfur compounds. Sulfur compounds rapidly deactivate both iron and cobalt catalysts, presumably by forming surface metal sulfides that do not have FTS activity. Ideally, there should be no sulfur in the syngas but there is always a small amount that gets through to the catalyst. There is really no safe sulfur level in FTS. Again, the level of gas cleaning required is based on economic considerations; namely how long the catalyst remains active versus the investment in gas cleaning. In the beginning, Fischer recommended 4 ppm as the maximum sulfur content in the FTS feedgas. Dry (1981) recommends a maximum sulfur content of 0.2 ppm based on commercial experience in the Sasol plants. Co catalysts are more sensitive to sulfur poisoning than Fe catalysts. Given the relative cost of Co versus Fe, more efficient sulfur removal should be expected for FTS with Co catalysts. Several more recent references cite even lower sulfur tolerances. Boerrigter, et al, (2002) state that sulfur levels in syngas for FTS should be below 1 ppmv and Turk, et al, (2001) claim that the total sulfur content in syngas should be 60 ppb. In general, for a fixed bed reactor design, any catalyst poison will have the most pronounced affect near the gas inlet and propagate through the reactor towards the outlet whereas in a fluidized bed design the poison will have a uniform affect throughout the reactor. Other syngas impurities are also known to poison FTS catalysts. Halide levels in syngas should be less than 10 ppb (Boerrigter, et al, 2002) and referenced nitrogen levels are 10 ppmv NH3, 0.1 ppmv NOx and 10 ppb HCN (Turk, et al, 2001). Additionally, water oxidizes FT catalysts (both Fe & Co) but the rate of oxidation is higher for the iron catalyst. Plus water has an inhibiting effect on the iron catalyst because of its WGS activity (Espinoza, et al, 2000). Table 23 summarizes the syngas impurities and tolerances.


Table 21: Syngas Impurities and Tolerances for Fisher-Tropsch

Sulfur Halides Nitrogen

Level 0.2 ppm 1 ppmv 60 ppb 10 ppb 10 ppmv NH3 0.1 ppmv NOx 10 ppb HCN

Source Dry, 1981 Boerrigter, et al, 2002 Turk, et al, 2001 Boerrigter, et al, 2002 Turk, et al, 2001

Note: There are differing points of view regarding sulfur level but in general, the sulfur content of the syngas should be minimized according to economics.

Commercial processes are available to clean syngas to meet these stringent contaminant requirements. The Rectisol process uses chilled methanol to scrub the raw syngas. NH3, H2S, tars, and CO2 are removed from syngas to required levels. Other chemical absorption processes include potassium carbonate or alkanolamine (MEA monoethanolamine or DEA diethanolamine) for wet scrubbing. Fixed bed reactors containing ZnO are also used for sulfur polishing. Whether or not these gas-cleaning processes are economical will depend on the scale of the FTS process.


Status/Technology Developers

The first FT plants began operation in Germany in 1938. There were nine plants with a total capacity of 660,000 tonnes/yr of FT products. These plants closed down after the second world war (Dry, 2002). Then in 1955, Sasol, a world-leader in the commercial production of liquid fuels and chemicals from coal and crude oil, started up its first FT plant (Sasol I) in Sasolburg, South Africa. The plants capacity is 6 million tonnes/yr of FT products from coal. There are five tubular fixed bed (ARGE) reactors for wax production that are still operational today and three circulating fluid bed (CFB) reactors (Synthol reactors), which operated until 1993 (Dry, 1982). In 1993, a slurry reactor that matches the production of the five ARGE reactors began operation. Following the success of Sasol I, Sasol II and III came on line in 1980 and 1982, respectively. These plants are located in Secunda, South Africa. Sasol II and III have all circulating fluid bed (CFB) reactors (Synthol reactors) but with improved heat exchange and thus they get a three-fold increase in capacity compared to the synthol reactors at Sasol I. From 1995 to 1999, these second generation CFB reactors (Synthol reactors;16 of them) were replaced with eight fixed fluid bed (FFB) reactors known as Sasol Advanced Synthol reactors (Dry, 2002). This was done for economic reasons and although the plant capacity did not increase the operating cost was projected to decrease by $1/BBL. In general, these reactors are advantageous due to lower capital costs, higher efficiency, lower operating costs, and less downtime for maintenance. CFB reactors suffered, initially, from enormous scale-up problems (mechanical, operational, & catalyst point of view). Sasols South Africa plants use both low and high temperature FT synthesis. The coal consumption at Sasols three commercial plants is 36 million tons/year that generates, on average, 1.5 million Nm3/hr of syngas from 97 pressurized Lurgi coal gasifiers (Van Nierop, et al, 2000). The syngas is converted to "more than 200 fuel and chemical products at its plants in Sasolburg and Secunda in South Africa as well as at several


other plants abroad." ( Its products; including gasoline, diesel, candle waxes, hard waxes, hydrocarbon lubricants, methane, phenol and cresol, tar and pitch, ammonia, detergents, etc., are exported to more than 70 countries around the world. Sasol also supplies 41% of South Africas liquid transportation fuel requirements. The company has developed world-leading technology for the commercial production of synthetic fuels and chemicals from lowgrade coal as well as the conversion of natural gas to environment-friendly diesel and chemicals. In the early 1990s, two other FT plants came on line. The Mossgas plant which converts natural gas to FT products using a high temperature process and an iron catalyst started up in South Africa in 1992 (Dry, 2002). This plant produces 1 million tonnes/yr of FT products including motor gasoline, distillates, kerosene, alcohols and LPG. Additionally, Shell commissioned a plant in 1993 in Bintuli, Malaysia using the Shell Middle Distillate Synthesis process, which is essentially enhanced FT synthesis. This plant produces 500,000 tonnes/yr (12,000 BPD) of FT products from 100 MMSCFD of natural gas using a cobalt catalyst (Senden, et al, 1992). The plant produces automotive fuels, specialty chemicals, and waxes. Also, Syntroleum is building a 10,000 BPD plant (specialty chemicals and lube oil) located in the Burrup Peninsula in Northwestern Australia (Haid, et al, 2000). The engineering began in 1999. Construction started in 2001 and operation is expected to begin in 2003. The plant will convert natural gas in their gas-to-liquids and specialty products plant using their proprietary process and a cobalt catalyst. In early 2000, Sasol studied the feasibility of replacing coal with recently discovered offshore natural gas. They decided that switching to natural gas would reduce expenditures in its coal mining operations as well as the high costs of compliance with environmental regulations associated with coal. Therefore, they are beginning to build a network of natural gas pipelines, which are expected to deliver gas to their facilities in 2004. Sasol plans to use the natural gas as a supplementary feedstock at Secunda and as a replacement feedstock for the coal in Sasolburg ( and


Process Performance

In reforming methane about 20% of the carbon is converted to CO2 while for coal this number is 50% due to the lower hydrogen content of coal (Dry, 2002). Coal gasification produces a syngas with a H2:CO ratio of about 0.67. For this ratio, the maximum obtainable conversion of CO to FT products without WGS is 33% (Raje, et al, 1997). The theoretical max conversion of methane to paraffins is 78% on a LHV basis. The conversion at the Bintuli plant in Malaysia is about 63% (Senden, et al, 1992). In a slurry reactor, very small catalyst particles are suspended in a liquid medium allowing effective heat removal and good temperature control. Sasols slurry phase process has a thermal efficiency of about 60% and a carbon conversion efficiency of about 75% (Lutz, 2001). The slurry reactor has several advantages over the multitubular design. For example, the capital cost is 25% less for the same capacity, the pressure difference across the


reactor is less and thus the gas compression cost is lower, there is a four fold lower catalyst consumption rate, and the reactor can operate at higher temperatures and thus higher conversions (Dry, 1982). One disadvantage is that any poison will poison the entire amount of catalyst whereas in the fixed tube design only the top portion of the catalyst will be poisoned.


Environmental Performance

Compared to conventional fuels, FT fuels contain no sulfur and low aromatics. These properties along with a high cetane number result in superior combustion characteristics (Jager, 1998 and Alleman and McCormick, 2003). Tests performed on heavy duty trucks showed decreases in vehicle emissions of HC, CO, NOx, and PM when using a FT fuel (Haid, et al, 2000; Lutz, 2001). FT diesel has been tested in a variety of lightand heavy-duty vehicles and engines. The paper by Alleman and McCormick (2003) summarizes FT diesel fuel property and emission information found in the literature. Overall, FT diesel showed a reduction in regulated as well as some unregulated emissions compared to conventional diesel. Several life cycle assessments (LCA) have been performed on a variety of transportation fuels including FT diesel and gasoline but most of the studies have examined only greenhouse gas emissions and energy consumption (General Motors, 2001; Marano and Cieferno, 2001; and MIT, 2000) with the exception of General Motors. They also examined five criteria pollutants (VOCs, CO, NOx, PM10, and SOx). These analyses have examined the emissions and energy consumption from resource extraction to end use. The results of the studies vary based on the feedstock procurement, technology conversion, and vehicle assumptions. However, in general, there is not a big advantage to FT liquids from fossil fuels in terms of energy consumption and green house gas emissions. This will not be the case for biomass systems. Because of the improved combustion characteristics of the FT liquids, performing a complete LCA including criteria pollutants will mostly likely show the overall benefits of FT liquids compared to conventional transportation fuels.

6.10 Cost
Table 24 gives information about the price of natural gas, gasoline, and diesel fuel, prior to giving cost information about FT synthesis. Over the years the price of each of these has fluctuated significantly.


Table 22: Price Ranges for Natural Gas, Gasoline, and Diesel
Average U.S. price (1996-2002) Natural gas $2.76/Mscf (national wellhead) $1.33/gallon (includes taxes) yearly avg $4.12/Mscf $3.8/GJ (2001) High Low Current price (Nov 02 avg) $3.64/Mscf $3.4/GJ

monthly avg $8.06/Mscf $7.4/GJ (Jan 01)

yearly avg $1.96/Mscf $1.8/GJ (1998)

monthly avg $2.35/Mscf $2.2/GJ (Jan 02)


$1.82/gallon $14/GJ (HHV); $15/GJ (LHV) (May 2001)

$0.97/gallon $7/GJ (HHV); $8/GJ (LHV) (Feb. 1999)

$1.42/gallon ($11/GJ, HHV; $12/GJ, LHV) Excluding dist., marketing & taxes: $0.74/gallon ($5.6/GJ, HHV; $6.1, LHV) $1.42/gallon ($9.7/GJ, HHV; $10.5, LHV) Excluding dist., marketing & taxes: $0.75/gallon ($5.1/GJ, HHV, $5.5, LHV)


$1.25/gallon (includes taxes)

$1.67/gallon $11/GJ (HHV); $12/GJ (LHV) (Oct. 2000)

$0.95/gallon $6/GJ (HHV); $7/GJ (LHV) (Feb. 1999)

Data from EIA: (1) U.S. Natural Gas Prices and Natural Gas Summary from the Short-Term Energy Outlook ( ) (Note: For natural gas, $/Mscf is essential equivalent to $/MMBtu) (2) Gasoline and Diesel Fuel Update ( ) Note: Mscf=1,000 standard cubic feet

The following is cost information specific to the existing commercial plants and studies by others.
(1) The capital cost of the Bintuli, Malaysia plant was about $660 million. (Senden, et al, 1992) (2) The capital cost of replacing Sasols 16 CFB reactors with 8 FFB reactors was $225 million. There was no increase in capacity but the reduction in the operating cost was projected to be $1/BBL. (Chang, 2000) (3) Studies by Sasol indicate that a two train GTL plant made up of slurry phase reactors producing 30,000 BPD of liquid transport fuels can be constructed at a capital cost of about $25,000/BPD capacity including utilities, offsite facilities, and infrastructure. (Lutz, 2001 and Vosloo, 2001) (4) Sasol examined the economics of a grass root plant using slurry phase reactors feeding 100 MMSCFD of natural gas and producing 10,000 BBL/day (425,000 ton/yr) of transportation fuel. The capital investment was around $300 million resulting in an investment of $30,000/daily BBL. (Jager, 1998) (5) At a natural gas price of $0.50/MMBtu (inexpensive), the feedstock accounts for $5/BBL of the product price and the operating cost (fixed and variable) is estimated


to contribute an additional $5/BBL to the product price. (Gradassi, 1998; Jager, 1998; Lutz, 2001; and Vosloo, 2001) (6) Using literature information, an economic analysis was performed for a 50,000 BPD plant using a feedstock cost of roughly $0.50/MMBtu (inexpensive). The resulting cost of the FT liquids at an internal rate of return of 15% was $26/BBL ($0.83/gallon). This means that in order for the natural gas to FT liquids process to be competitive at this low feedstock cost, the price of crude oil needs to be at least $18/BBL. (Gradassi, 1998) (7) FT liquids from both coal and natural gas were examined for a plant that produces roughly 50,000 BPD of FT liquids. At a natural gas cost of $0.5/MMBtu (inexpensive) the cost of the FT liquids is $24/BBL ($0.76/gallon) and at $4/MMBtu this cost increases to $52/BBL ($1.7/gallon). Using Illionois #6 coal, the cost of the FT liquids is $46/BBL ($1.5/gallon). For a once-through plant that co-produces electricity which is sold at $0.05/kWh, the cost of the FT is reduced to $35/BBL ($1.1/gallon). (Gray and Tomlinson, 1997) The majority of the capital cost, greater than 50%, comes from syngas generation (Dry 2002; Senden, et al, 1992, and Vosloo, 2001). Additionally, the cost of syngas production from methane is 30% lower than that from coal and it is a more efficient process (Dry 2002). Any cost reduction in syngas generation will have a large impact on the economics. Three studies that have examined the production of FT liquids from biomass are summarized in Table 25.


Table 23: Comparison of Biomass Based Fischer Tropsch Studies Study Biomass feed rate (BD tonne/day) Biomass cost Electricity selling price Net power (MW) Other fuels Raw or finished products FT fuels produced Cost year Price of products (without distribution costs and taxes) Level of detail Tijmensen (2000) (a), (b) 1,920 $2/GJ $38/dry tonne $0.057/kWh 80 100 from biomass (150 MW gas turbine) co-fires nat gas in gas turbine raw 1,216 - 2,026 BBL/day 1999 $13 - $30/GJ $1.8 - $4.1/gallon base case = $15 - 16/GJ Mitre (1996) (a), (b) 2,000 $2.45/GJ $46/dry tonne $0.05/kWh 110 - 120 none finished 1,357 - 1,715 BBL/day 1993 $8 - $14/GJ $1.1 - $1.9/gallon Novem (2000) (a), (c) 1,358 $3/GJ $55/dry tonne $0.067/kWh about 150 co-fires nat gas in gas turbine finished 486 - 1,378 BBL/day 1999 $9 - $13/GJ $1.2 - $1.8/gallon

Very detailed - Gives Somewhat detailed Some details - Gives costs of individual Gives costs of major costs of major plant equipment, operating plant sections, sections, and costs, and other operating costs, and assumptions used to economic parameters lists other economic determine installed and assumptions assumptions costs Notes: (a) Examined direct & indirect gasifiers and atmospheric and pressurized gasifiers. (b) Examined both maximum liquids (recycle of unconverted syngas) and once through with power generated from unconverted syngas. (c) Only examined once through with combined heat and power.


6.11 References
Abbott, J. and Crewdson, B. (2002). "Gas heated reforming improves Fischer-Tropsch process." Oil & Gas Journal 100(16): 64-66, 68, 70. Adesina, A.A. (1996). "Hydrocarbon synthesis via Fischer-Tropsch reaction: travails and triumphs", Appl. Catal. A, 13: 345-367. Alleman, T.L. and McCormick, R.L. (2003). "Fischer-Tropsch Diesel Fuels - Properties and Exhaust Emissions: A Literature Review". 2003 SAE World Congress, CI Engine Combustion Processes & Performance with Alternative Fuels (SP-1737), Detroit, Michigan, March 3-6. Bartholomew, C. H. (1991). "Recent Developments in Fischer-Tropsch Catalysis," Chapter 5 in "New Trends in CO Activation", (L. Guczi ed.), Stud. Surf. Sci. Catal., 64: 158. Boerrigter, H., den Uil, H., and Calis, H.-P. (2002). "Green Diesel from Biomass via Fischer-Tropsch Synthesis: New Insights in Gas Cleaning and Process Design." Paper presented a Pyrolysis and Gasification of Biomass and waste, Expert Meeting, 30 September, 2002, Strasbourg, FR. Chang, T. (2000). "South African company commercializes new F-T process." Oil & Gas Journal 98(2): 42-45. Davis, B. H. (2001). "Fischer-Tropsch synthesis: current mechanism and futuristic needs," Fuel Proc. Tech., 71(1-3): 157-166. Dry, M. E. (2002). "The Fischer-Tropsch process: 1950-2000." Catalysis Today 71(3-4): 227-241. Dry, M. E. (1999). "Fischer-Tropsch reactions and the environment", Appl. Catal. A, 189(2): 185-190. Dry, M.E. (1996). "Practical and Theoretical Aspects of the Catalytic Fischer-Tropsch Process," Appl. Catal. A, 138: 319 Dry, M. E. (1990). "The Fischer-Tropsch process-commercial aspects", Cat. Today, 6: 183. Dry, M. E. (1988). "The Sasol route to chemicals and fuels." Studies in Surface Science and Catalysis 36(Methane Convers.): 447-56. Dry, M. E. (1982). "Catalytic Aspects of Industrial Fischer-Tropsch Sunthesis." Journal of Molecular Catalysis 71: 133-144. Dry, M.E. (1981). "The Fischer-Tropsch Synthesis," in Catalysis, Science and Technology, (J.R. Anderson, M. Boudart, eds.), Springer-Verlag, Vol. 1, pp. 159256.


Dry, M. E. and Hoogendoorn, J.C. (1981). "Technology of the Fischer-Tropsch process." Catal. Rev. - Sci. Eng. 23(2): 265-78. Espinoza, R. L., Du Toit, E., Santamaria, J., Menendez, M., Coronas, J. and Irusta, S. (2000). "Use of membranes in Fischer-Tropsch reactors." Studies in Surface Science and Catalysis 130A(International Congress on Catalysis, 2000, Pt. A): 389-394. Espinoza, R. L., Steynberg, A. P., Jager, B., and Vosloo, A. C. (1999). "Low temperature Fischer-Tropsch synthesis from a Sasol perspective", Appl. Catal. A, 186(1-2): 13-26. Frohning, C., Kolbel, H., Ralek, M., Rottig, W., Schnur, F., and Schulz, H. (1982). "Fischer-tropsch process", Chapter 8 in Chemical feedstocks from coal, J. Falbe, ed., John Wiley and Sons, NY, pp. 309-432. General Motors, Argonne National Laboratory, ExxonMobil, Shell. (2001). "Well-toWheel energy Use and Greenhouse Gas emissions of Advanced Fuel/Vechile Systems - North American Analysis." Gradassi, M. J. (1998). "Economics of gas to liquids manufacture." Studies in Surface Science and Catalysis 119(Natural Gas Conversion V): 35-44. Gray, D. and Tomlinson, G. (1997). "A technical and economic comparison of natural gas and coal feedstocks for Fischer-Tropsch synthesis." Studies in Surface Science and Catalysis 107(Natural Gas Conversion IV): 145-150. Haid, M. O., Schubert, P. F. and Bayens, C. A. (2000). "Synthetic fuel and lubricants production using gas-to-liquids technology." DGMK Tagungsbericht 2000-3: 205212. Jager, B. (1998). "Developments in Fischer-Tropsch technology." Studies in Surface Science and Catalysis 119(Natural Gas Conversion V): 25-34. Jager, B., and Espinoza, R. L. (1995). "Advances in low temperature Fischer-Tropsch synthesis", Cat. Today, 23: 17. Jess, A., Popp, R., and Hedden, K. (1999). "Fischer-Tropsch-synthesis with nitrogen-rich syngas: Fundamentals and reactor design aspects", Appl. Catal. A, 186(1-2): 321-342. Kolbel, H., Ralek, M. (1980). "Fischer-Tropsch synthesis in the liquid phase", Catal. Rev. - Sci. Eng., 21(2): 225-274. Lutz, B. (2001). "New age gas-to-liquids processing." Hydrocarbon Engineering 6(11): 23-26,28.

Marano, J. J., and Ciferno, J. P. (2001). "Life-Cycle Greenhouse-Gas Emissions Inventory For Fischer-Tropsch Fuels". U.S. DOE. National Energy Technology Laboratroy. Massachusetts Institute of Technology Energy Laboratory. (2000). "On the Road in 2020: A Life-Cycle Analysis of New Automobile Technologies". MIT EL 00-003. Mills, G. (1993) Status and future Opportunities for Conversion of Synthesis Gas to Liquid Energy Fuels: Final Report. National Renewable Energy Laboratory, Golden, CO, TP-421-5150. MITRE Corporation. (1996). Techno-Economic Assessment of Biomass Gasification Technologies for Fuels and Power. Novem. (2000). "Technical and Economic Data Biomass-Bases Energy Conversion Systems for the Production of Gaseous and/or Liquid Energy Carriers". Report No. GAVE-9915. Oukaci, R., Singleton, A. H., and Goodwin, J. G. (1999). "Comparison of Patented Co FT Catalysts using Fixed-Bed and Slurry Bubble Column Reactors", Appl. Catal. A, 186: 129-144. Patzlaff, J., Liu, Y., Graffmann, C., and Gaube, J. (1999). "Studies on product distributions of iron and cobalt catalyzed Fischer-Tropsch synthesis", Appl. Catal. A, 186(1-2): 109-119. Raje, A., Inga, J. R. and Davis, B. H. (1997). "Fischer-Tropsch synthesis: process considerations based on performance of iron-based catalysts." Fuel 76(3): 273280. Rao, V. U. S, Stiegel, G. J., Cinquegrane, G. J., and Srivastave, R. D. (1992). "Iron based catalysts for slurry phase Fischer-Tropsch process: Technology Review", Fuel Proc. Tech., 30: 83-107. Rostrup-Nielsen, J. R. (2000). "New Aspects of Syngas Production and Use." Catalysis Today 63: 159-167. Schulz, H. (1999). "Short history and present trends of Fischer-Tropsch synthesis", Appl. Catal. A, 186(1-2): 3-12. Schulz, H. (1985). "Selectivity and mechanism of the Fischer-Tropsch COhydrogenation", C1 Mol. Chem., 1: 231 Senden, M. M. G., Sie, S. T., Post, M. F. M. and Ansorge, J. (1992). "Engineering aspects of the conversion of natural gas into middle distillates." NATO ASI Series, Series E: Applied Sciences 225(Chemical Reactor Technology for Environmentally Safe Reactors and Products): 227-47.


Steynberg, A. P., Espinoza, R. L., Jager, B., and Vosloo, A. C. (1999). "High temperature Fischer-Tropsch synthesis in commercial practice", Appl. Catal. A, 186(1-2): 41-54. Tijmensen, M. J. A., Faaij, A. P. C., Hamelinck, C. N. and van Hardeveld, M. R. M. (2002). "Exploration of the possibilities for production of Fischer Tropsch liquids and power via biomass gasification." Biomass & Bioenergy 23(2): 129-152. Tijmensen, M. J. A., Faaij, A. P. C., Hamelinck, C. N. and van Hardeveld, M. R. M. (2000) "The production of Fischer Tropsch liquids and power through biomass gasification", M. Sc. Thesis, Utrecht University/Science, Technology and Society, Utrecht. Report No. NWS-E-2000-29. Turk, B.S., Merkel, T., Lopez-Ortiz, A., Gupta, R.P. Portzer, J.W., Kishnam, G., Freeman, B.D., and Fleming, G.K. (2001). "Novel TEchnologies for Gaseous Contaminants Control." Final Reprot for DOE Contract No. DE-AC2699FT40675, September 2001. Van Nierop, P., Erasmus, H.H., and van Zyl, J.J., (2000). "Sasol's Achievements if the 20th Century as a Building Block for the 21st." Presented at the 2000 Gasification Technologies Council, October 8-11, 2000. San Francisco, CA Vannice, M. A. (1976). "The Catalytic Synthesis of Hydrocarbons from Carbon Monoxide and Hydrogen", Catal. Rev.-Sci. Eng., 14(2): 153-191. Wender, I. (1996). "Reactions of synthesis gas." Fuel Processing Technology 48(3): 189-297.


7.1 Introduction

The oxysynthesis process is the hydroformylation of olefins with synthesis gas (CO and H2) and is the principal chemical route to C3-C15 aldehydes that are typically converted into alcohols, acids, or other derivatives. Worldwide production of oxo-aldehydes and alcohols was 6.5 million tons per annum in 1997 (Bahrmann and Bach 2000). The announced planned capacity additions for oxo chemicals in the next 5 years is almost an additional 1 million tons (Bitzzari et al. 2002). Hydroformylation is an industrial synthetic route for the conversion of a large number of olefins and is used to produce solvents, synthetic detergents, flavorings, perfumes and other healthcare products, and other high value commodity chemicals. Oxo aldehyde products range from C3-C15. The higher end of this range (C11-C14) is often referred to as the detergent range because these aldehydes are typically used as intermediates in the synthesis of alcohols with the desired surfactancy and lack of crystallinity appropriate for detergents. The hydroformylation of olefins was discovered in 1938 by Roelen of Ruhrchemie as an outgrowth of Fischer-Tropsch Synthesis (FTS). Oxosynthesis conducted in Germany was thought to be a heterogeneously catalyzed reaction because of its association with FTS. Continuous addition of cobalt, however, was required for early oxosynthesis processes. This prompted additional investigation that identified soluble cobalt carbonyls as the active catalyst for the reaction making the oxosynthesis process one of the first significant homogeneously catalyzed industrial processes. Reactants, catalyst, and products are all present in the same phase, usually as liquids. Today, hydroformylation processes are the fourth largest commercial use of syngas (Wender 1996). Cobalt hydrocarbonyl was the first oxosynthesis catalyst commercially used, however, in the mid-1970s more active and selective ligand-modified rhodium-based catalysts were introduced for hydroformylation of lower olefins. For example, n-butanal from propylene is the highest volume oxo product. It can be converted to n-butanols (used as solvents) or 2-ethyhexanol, a plasticizer alcohol.


Technology Description

Oxosynthesis processes are very versatile and are used to commercially produce a wide range of products. The highest volume oxosynthesis process is the hydroformylation of propylene to butyraldehyde, which in turn is converted to 2-ethylhexanol, a plasticizer alcohol used to make flexible PVC. Hydroformylation is a key step in the commercial production of Vitamin A. Hydroformylation of allyl alcohol is used to commercially produce butanediol.



The oxosynthesis process is also known as hydroformylation and involves the reaction of CO and H2 with olefinic hydrocarbons to form an isomeric mixture of normal- and iso86

aldehydes. The basic oxosynthesis reaction is highly exothermic and is thermodynamically favorable at ambient pressures and low temperatures (Whyman 1985). The reaction proceeds only in the presence of homogeneous metal carbonyl catalysts. One of the more important factors in oxosynthesis is the normal to branched isomeric ratio (n/i). The normal, or straight chain isomer is the desired product. RCH=CH2 + CO + H2 RCH2CH2CHO (normal) + R(CH3)CHCHO (branched) Usually a 1:1 H2:CO syngas mixture is required for oxosynthesis. The overall reaction rate has first-order dependence on the hydrogen partial pressure and inverse-first order dependence on CO partial pressure making the reaction rate essentially independent of total pressure (Pruett 1979). Higher CO partials pressures are usually required, however, to maintain the stability of the metal carbonyl catalysts. The reaction is also first order in olefin and metal concentration at the higher CO partial pressures. The mechanism for the homogeneously catalyzed hydroformylation reaction is a function of the catalyst system used in the process. However, the general steps in the reaction paths are similar. The first step is the removal of CO from the catalyst organometallic complex to give a coordinatively electron deficient species. The olefin attaches to the vacated d-orbital at this site in the catalyst complex. The double bond in the olefin attaches to the metal atom yielding an alkyl metal carbonyl complex. The olefin inserts into the M-H bond. CO then inserts into the complex at the C-M bond followed by hydrogen insertion at the same point to yield the aldehyde product. A generalized reaction mechanism taken from Whyman (1985) follows:

M 2 (CO ) 6 L2 + H 2 2 HM (CO ) 3 L 2 HM (CO ) 2 L + 2CO RCH =CH 2 + HM (CO ) x L H ( RCH =CH 2 ) M (CO ) x L H ( RCH =CH 2 ) M (CO ) x L RCH 2 CH 2 M (CO ) x L RCH 2 CH 2 M (CO ) x L + CO RCH 2 CH 2 COM (CO ) x L RCH 2 CH 2 COM (CO ) x L + H 2 RCH 2 CH 2 CHO + HM (CO ) x L RCH 2 CH 2 COM (CO ) x L + HM (CO ) 3 L RCH 2 CH 2 CHO + M 2 (CO ) 3+ x L2
M= metal (Co or Rh) L = ligand (CO, PR3, or PPh3) x = 2 or 3



Oxosynthesis is a rapid reaction catalyzed by soluble cobalt or rhodium complexes. The advances made in organometallic chemistry in the 1950s and 1960s led to improved hydroformylation synthesis catalysts with higher thermal stability and greater selectivity. Three complimentary catalytic hydroformylation processes have been developed and commercialized. The choice of catalyst depends of the particular starting olefin or desired product. Phosphine-modified rhodium catalysts are best to produce 2ethylhexanol from propylene and syngas. Unmodified cobalt catalysts are versatile enough to be used for hydroformylation of high carbon number mixed olefins. Phosphine-modified cobalt catalysts are used for the production of higher, detergent


range, alcohols rather than aldehydes. discussed.

Each of these catalyst systems will be

7.4.1 Cobalt Carbonyl

The first hydroformylation catalysts were cobalt carbonyls, the specific active catalyst was found to be HCo(CO)4 in equilibrium with Co2(CO)8 based on the following reaction: Co2(CO)8 + H2 2 HCo(CO)4 Cobalt metals and most cobalt salts will form cobalt carbonyl under hydroformylation conditions. The cobalt catalyzes both double bond isomerization and oxosynthesis. Undesired competing side reactions such as the direct hydrogenation of the starting olefin and the condensation of product aldehydes to high boiling products are generally avoided in the Co-catalyzed process. For cobalt carbonyl catalysts, a n/i ratio of 4:1 can be achieved with catalyst concentrations of 0.1-1% metal/olefin at 200-300 atm and 110-200C with a 1:1 H2/CO ratio. Lower process temperatures and higher CO partial pressures favor the formation of the straight chain isomer, however, the overall conversion efficiency decreases. The cobalt carbonyl catalysts are also not very stable at the higher process temperatures. Catalyst tends to deposit on reactor walls decreasing catalyst activity and reducing catalyst recovery at elevated process temperatures.

7.4.2 Phosphine-modified Cobalt

In the early 1960s Shell Oil Company commercialized a new cobalt-based hydroformylation process. The addition of a phosphine ligand to Co resulted in a trialkylphosphine-substituted cobalt carbonyl catalyst [HCo(CO)3P(n-C4H9)3]. This new catalyst had higher selectivity to straight-chain aldehydes (n/i = 7:1) with improved thermal stability compared to the unsubstituted cobalt catalysts. The improved thermal stability allows for lower process pressures but higher process temperatures. The Shell process conditions are 50-100 atm and 160-200C with H2:CO = 1. Even though this catalyst has improved thermal stability it has a lower hydroformylation activity than cobalt carbonyl catalysts, hence the higher reaction temperature. The higher temperatures also increase the competing olefin hydrogenation reaction. Shell has optimized this process to produce detergent range alcohols (C11-C14) in a single step by capitalizing on the conversion of terminal olefins to alcohols by hydrogenation of the aldehyde hydroformylation products. A high n/i ratio results from increased isomerization rates concurrently with hydroformylation with the phosphinemodified cobalt catalyst.

7.4.3 Phosphine-modified Rhodium

Significantly lower operating pressures and temperatures and increased selectivity to linear products was demonstrated with the introduction of phosphine-modified rhodium catalysts. In 1976, Union Carbide and Davy Process Technology commercialized the LP Oxo process based on a triphenylphosphine-modified rhodium catalyst. With a HRh(CO)(PPh3)3 catalyst the process is maintained at 7-25 atm and 60-120C and a n/i


ratio of 8-12:1 is possible. Low temperatures, high carbon monoxide partial pressure, and high ligand concentration on the Rh catalysts favors the formation of the straightchain isomers. Rh-based catalysts are mainly used for the hydroformylation of lower olefins (propylene to butyraldehyde, for example) but have limited use for higher olefins because of thermal instability of the catalyst at the high temperatures required for distillation to separate product and catalyst. Another important factor when selecting Rhbased hydroformylation catalysts is the high cost and low availability of rhodium. This makes catalyst separation and recycle an economically important process consideration. The high cost of rhodium, however, is offset by lower equipment costs, increased activity, and higher selectivity and efficiency. The development of water-soluble Rhbased catalysts avoids some of these issues. Rhone-Poulenc commercialized an oxo process based on a water-soluble Rh catalyst in 1984 (Billig and Bryant 2000).


Gas Cleanliness Requirements

Catalyst lifetimes are significantly reduced by poisoning from strong acids, HCN, organosulfur, H2S, COS, O2, and dienes (Bahrmann and Bach 2000).


Status/Technology Developers

The LP Oxo Low Pressure Oxo Process is the result of a collaboration between Davy Process Technology (DPT) and Union Carbide Corporation that began in 1971 and is the world's leading process for the production of oxo alcohols from olefins. It represents most of the world's oxo alcohols capacity licensed during the last 20 years. A schematic of the LP Oxo Process is shown in Figure 10.
Compressor Off gas

Stirred-Tank Reactor Demister Condenser Stripping Column Reprocessing Purify Rh complex and ligand Offline Catalyst Reprocessing Separator

CO/H2 H2 Olefin



Figure 9: Schematic Flow Diagram of the LP Oxo Process


The LP Oxo Process involves the reaction of propylene with synthesis gas to produce butyraldehydes (n/i = 10-20:1), which are in turn converted to 2-ethylhexanol. Normal or iso-butyraldehydes are also converted to normal and iso-butanols and other derivatives. Propylene to mixed butyraldehyde efficiencies as high as 97.5% have been demonstrated at commercial scale. The efficiency is a function of feedstock purity with a higher purity feed resulting in a higher efficiency due to lower purge losses. The Shell Process is another commercial oxo synthesis process based on a phosphine modified cobalt catalyst to produce alcohols directly from olefins. The conversion of product oxo aldehydes to alcohols occurs in a single step in the Shell process because of the high hydrogenation activity of the modified catalyst. This process yields detergent range alcohols from a range of higher olefins. A summary of oxochemical producers in the U.S. as cited in the Chemical Economics Handbook (CEH) (Bitzzari et al. 2002) is presented in Table 26.


Table 24: Summary of U.S. Oxo Chemical Producers Company BASF Corporation, Chemicals Division Industrial Organics Celanese Ltd. Chemicals Division Dow Chemical Dow Chemical Plant Location Freeport, TX Bay City, TX Taft, LA Texas City, TX Annual Capacity as of 9/01/02 (tonnes x103/yr) 270 30 45 250 27 272 15 100 45 335 33 118 424 184 408 340 102 140 12

Oxo Products
n-butyraldehyde isobutyraldehyde propionaldehyde n-butyraldehyde iosbutyraldehyde n-butyraldehyde isobutyraldehyde propionaldehyde valeraldehyde n-butyraldehyde isobutyraldehyde propionaldehyde n-butyraldehyde

Rh Rh Rh

Final Products
n-butanol 2-ethylhexanol neopentyl glycol; isobutanol n-propanol n-butanol isobutanol n-butanol iso-butanol n-propanol; propionic acid valeric acid; amyl alcohol n-butanol; 2-ethylhexanol; 2-ethylhexanoic acid isobutanol n-propanol; propionic acid n-butanol; 2-ethylhexanol, 2-ethylheanoic acid; n-butyric acid; methyl amyl ketone neopentyl glycol; methyl isoamyl ketone; isobutyl acetate phthalate and trimellitate plasticizers Neodol alcohols Phthalate plasticizers 2-ethylhexanol sold without conversion


Eastman Chemical Company, Eastman Division ExxonMobil Chemical Company Shell Chemical Company Sterling Chemicals, Inc. Sunoco, Inc. From: (Bitzzari et al. 2002)

Longview, TX

Rh Baton Rouge, LA Geismar, LA Texas City, TX Pasadena, TX isobutyraldehyde Alcohols: C6-C13 branched C9-C11 linear C9-C15 linear alcohols C7, C9, C11 linear alcohols n-butyraldehyde isobutyraldehyde

Co Co-PR3 Co Rh




The cost of oxo products are dependent on the initial costs of the olefin feedstocks, primarily propylene and ethylene. The oxo aldeyhdes are effectively intermediates for the production of alcohols, acids and other chemical products. The market price for these products also varies with the cost of the feedstocks and intermediates. Chemical Economics Handbook (Bitzzari et al. 2002) contains estimated costs for a variety of oxo end products. The costs for the major products listed in Table 1 are presented in Table 2. Although not shown in Table 2, the highest priced oxo chemical in the CEH list is pelargonic acid ($1.33/lb). The majority of the oxo products listed in CEH fall in the range of 50-70/lb.
Table 25: U.S. Prices for Oxo Chemicals Product 2002 Price (/lb) n-butyraldehyde 52 n-butanol 55 2-ethylhexanol 63 2-ethyl hexanoic acid 68 isobutyraldehyde 52 Isobutanol 50 n-propanol 55 propionic acid 44 amyl alcohol 63 From: (Bitzzari et al. 2002)




Bahrmann, H., and Bach, H. (2000). "Oxo Synthesis." Ullman's Encyclopedia of Industrial Chemistry, Wiley - VCH Verlag GmbH & Co. Billig, E., and Bryant, D. R. (2000). "Oxo Process." Kirk-Othmer Encyclopedia of Chemical Technology, John Wiley & Sons. Bitzzari, S. N., Gubler, R., and Kishi, A. (2002). "Oxo Chemicals." Chemical Economics Handbook, SRI International, 1-121. Report number 682.7000. Pruett, R. L. (1979). "Hydroformylation." Advances in Organometallic Chemistry, 17, 160. Wender, I. (1996). "Reactions of synthesis gas." Fuel Processing Technology, 48(3), 189-297. Whyman, R. (1985). "Industrial Applications of Homogeneous Catalysts." Selected Developments in Catalysis, J. R. Jennings, ed., Blackwell Scientific Publications, Oxford, 128.


8.1 Introduction

The isosynthesis reaction is one that converts syngas over a thorium or zirconium-based catalyst at relatively extreme conditions to isobutene and isobutane. This reaction has often been considered a variation of Fischer-Tropsch Synthesis, however, there are major differences between the two synthetic processes (Feng et al. 1994). Isosynthesis is selective to i-C4 hydrocarbons whereas FTS forms a range of olefins according to the ASF distribution. Only trace amounts of oxygenates (water, methanol, isobutanol, DME, etc.) are formed under isosynthesis reaction conditions. Selective formation of branched hydrocarbons also occurs in isosynthesis. The isosynthesis reaction was first studied by Pilcher and Ziesecke in the early 1940s using ThO2-based catalysts and severe reaction conditions 150-1000 atm and 450C (Sofianos 1992). The goal was to optimize isobutane production for use in high-octane gasoline. The development of the petroleum industry led to waning interest in commercial isosynthesis processes and since these initial studies, isosynthesis has not received a lot of industrial attention. Demand for MTBE as a gasoline additive increased in the early 1990s to the point that supply shortages of petrochemical isobutene were becoming a concern (Sofianos 1992). This led to renewed interest in the isosynthesis reaction to convert syngas to branched hydrocarbons, particularly isobutene, which is a reactant in the synthesis of MTBE. Catalyst development has been focused on ZrO2-based catalysts for increased activity and selectivity to i-C4 products.


Technology Description

No commercial isosynthesis processes have been developed. Most laboratory studies have been conducted in gas-solid fixed bed reactors. Slurry reactors are also being investigated for isosynthesis, in fact, selectivities to C4 products are reported to be higher in slurry reactors compared to fixed bed reactors (Erkey et al. 1995). There is always a large amount of CO2 produced in the isosynthesis reaction, likely by the reaction of H2O and CO in the water-gas shift (Su et al. 2000a; Su et al. 2000b). CO2 recycle and reuse should be investigated to improve conversion efficiency to products.



Isosynthesis conditions are optimized for isobutene/isobutane production. At lower temperatures, alcohols and other oxygentates are formed. At higher temperatures, methane and aromatics are formed. The isosynthesis mechanism is related to that of higher alcohol synthesis. It involves 2 chain growth mechanisms - a step-wise CO insertion reaction and a condensation reaction mechanism involving surface adsorbed oxygenates. The precursor to the

oxygenate products has been identified as a surface adsorbed methoxide species (Maruya et al. 1996). The oxygenates, alcohols and ethers, are more than likely the primary reaction products that undergo dehydration and hydrogenation to form the isoalkenes and branched alkanes. These two competing chain growth mechanisms result in a discontinuity in the ASF distribution at C4 that explains the relatively high selectivity of the isosynthesis reaction to C4 products.



The first catalysts used by Pilcher and Ziesecke were ThO2-based. Unpromoted thoria catalysts have demonstrated 46% CO conversion at 150 atm and 450C but only 10% isobutene in the C4 fraction that is mostly isobutane (Sofianos 1992). DME was the main product at higher pressures. Thorium-based catalysts are good alcohol dehydration catalysts and in turn, are the most active isosynthesis catalysts. They have long lifetimes because they can be regenerated by oxidizing the accumulated coke that deposits on the surface and they are not poisoned by sulfur. Unfortunately, they are radioactive precluding their commercial use. Zirconium-based catalysts also have high activity for isosynthesis. Unpromoted zirconium catatalysts have demonstrated 32% CO conversion at 150 atm and 450C with much higher selectivity to isobutene (Sofianos 1992) compared to the thoriumbased catalysts. The overall activity of ZrO2-based catalysts for isosynthesis is lower than ThO2-based catalysts. Various promoters have been investigated to improve the activity and selectivity of ThO2 and ZrO2 catalysts (Jackson and Ekerdt 1990a; Jackson and Ekerdt 1990b). The most active isosynthesis catalyst is 20% Al2O3/ThO2. Other promoters such as Zn, Cr and alkali metals have also been tested. The addition of alkali metals to zirconium catalysts had a negative effect on catalyst performance (Jackson and Ekerdt 1990b; Li et al. 2001a). Doped zirconium-based catalysts posses oxygen vacancies in the oxide lattice. The most active catalysts tend to have maximum ionic conductivity suggesting that vacancies in the crystal lattice play an important role in the isosynthesis reaction (Li et al. 2001b; Li et al. 2002). These oxygen vacancy sites are required for methoxide formation on the catalyst surface that contributes to the condensation reaction (Jackson and Ekerdt 1990a). The selectivity of the isosynthesis reaction depends on the nature of the active catalyst sites, including oxygen vacancies on the surface, and the number of acidic and basic sites. The balance between acidic and basic catalyst sites dictates overall activity and selectivity (Li et al. 2002). Enhancing the number of acidic sites on the ZrO2 catalyst increased activity and selectivity to linear C4 hydrocarbons. Increasing the number of basic sites on the catalyst increased the yield of iso-C4 hydrocarbons (Li et al. 2001b). The acidic catalyst sites are thought to promote condensation and dehydration reactions. The basic sites are known to catalyze CO insertion reactions. The activity of promoted isosynthesis catalyst systems is related to the acid/base ratio, which can be altered and controlled by varying the preparation procedure for mixed oxide catalysts (Feng et al. 1995; Feng et al. 1994).



Gas Cleanliness Requirements

Thoria catalysts are not poisoned by sulfur and have high resistance to other poisons as well.


Status/Technology Developers

The isosynthesis process is not currently commercially practiced.



Erkey, C., Wang, J. H., Postula, W., Feng, Z. T., Philip, C. V., Akgerman, A., and Anthony, R. G. (1995). "Isobutylene Production from Synthesis Gas over Zirconia in a Slurry Reactor." Industrial & Engineering Chemistry Research, 34(4), 10211026. Feng, Z. T., Postula, W. S., Akgerman, A., and Anthony, R. G. (1995). "Characterization of Zirconia-Based Catalysts Prepared by Precipitation, Calcination, and Modified Sol-Gel Methods." Industrial & Engineering Chemistry Research, 34(1), 78-82. Feng, Z. T., Postula, W. S., Erkey, C., Philip, C. V., Akgerman, A., and Anthong, R. G. (1994). "Selective Formation of Isobutane and Isobutene from Synthesis Gas over Zirconia Catalysts Prepared by a Modified Sol-Gel Method." Journal of Catalysis, 148(1), 84-90. Jackson, N. B., and Ekerdt, J. G. (1990a). "Isotope Studies of the Effect of Acid Sites on the Reactions of C-3 Intermediates During Isosynthesis over Zirconium Dioxide and Modified Zirconium Dioxide." Journal of Catalysis, 126(1), 46-56. Jackson, N. B., and Ekerdt, J. G. (1990b). "The Surface Characteristics Required for Isosynthesis over Zirconium Dioxide and Modified Zirconium Dioxide." Journal of Catalysis, 126(1), 31-45. Li, Y. W., He, D. H., Cheng, Z. X., Su, C. L., Li, J. R., and Zhu, Q. M. (2001a). "Effect of calcium salts on isosynthesis over ZrO2 catalysts." Journal of Molecular Catalysis a-Chemical, 175(1-2), 267-275. Li, Y. W., He, D. H., Yuan, Y. B., Cheng, Z. X., and Zhu, Q. M. (2001b). "Selective formation of isobutene from CO hydrogenation over zirconium dioxide based catalysts." Energy & Fuels, 15(6), 1434-1440. Li, Y. W., He, D. H., Yuan, Y. B., Cheng, Z. X., and Zhu, Q. M. (2002). "Influence of acidic and basic properties of ZrO2 based catalysts on isosynthesis." Fuel, 81(11-12), 1611-1617.


Maruya, K., Takasawa, A., Haraoka, T., Domen, K., and Onishi, T. (1996). "Role of methoxide species in isobutene formation from CO and H- 2 over oxide catalysts - Methoxide species in isobutene formation." Journal of Molecular Catalysis aChemical, 112(1), 143-151. Sofianos, A. (1992). "Production of Branched-Chain Hydrocarbons Via Isosynthesis." Catalysis Today, 15(1), 149-175. Su, C. L., He, D. H., Li, J. R., Chen, Z. X., and Zhu, Q. M. (2000a). "Influences of preparation parameters on the structural and catalytic performance of zirconia in isosynthesis." Journal of Molecular Catalysis a-Chemical, 153(1-2), 139-146. Su, C. L., Li, J. R., He, D. H., Cheng, Z. X., and Zhu, Q. M. (2000b). "Synthesis of isobutene from synthesis gas over nanosize zirconia catalysts." Applied Catalysis a-General, 202(1), 81-89.