The Direct Methanol Fuel Cell

In the 1960s, the hope for a "direct" fuel cell emerged again, but instead of using coal, the fuel was methanol. Methanol, as opposed to coal, could be electrochemically oxidized at the electrodes. At the time, methanol was being used in fuel cells "indirectly," undergoing steam reforming to produce hydrogen, but, if a fuel cell could use methanol directly to produce electricity, then the fuel reforming step could be circumvented, allowing for a simpler system.

Before the major efforts of the 1960s, there was an instance of the use of methanol by Kordesch and Marko (1951). They described the development and performance of new carbon electrodes and identified "new possibilities for the building of fuel cells," such as the use of aldehyde (formaldehyde) and alcohols (ethanol and methanol) as fuel. The electrical current from alcohol was lower than that of formaldehyde, which could supply a small current of 0.3 mA/cm2 at 0.8 or 0.9 V using carbon electrodes in an alkaline electrolyte of KOH.

Direct methanol fuel cells were developed by researchers at Shell (Williams et al., 1965) and ESSO (Tarmy and Ciprios, 1965) with aqueous acid electrolytes that would not react with the CO2 produced in the electrochemical reaction. Alkaline electrolytes were tested by researchers at Allis-Chalmers (Murray and Grimes, 1963), who expected the degradation of the electrolyte by carbonate formation but also recognized the better compatibility of materials. For methanol oxidation on the anode, catalysts based on alloys between noble metals were more effective than pure metal catalysts, and Binder et al. (1965) studied different combinations in both acid and alkaline electrolytes. In 1992, research was revived following the technical improvements of the solid polymer fuel cell, as scientists at the Jet Propulsion Laboratory developed a direct methanol fuel cell using the same solid polymer electrolyte (Surampudi et al., 1994).

2.6.1 Acid Electrolyte: Shell (1965) and ESSO (1965)

At Shell Research Limited (Thornton Research Centre), Williams et al. (1965) chose acid over alkaline as the electrolyte because it was unaffected by the carbon dioxide produced in the methanol oxidation reaction. Also, with acid electrolytes, the water removal from the fuel cell was simpler. Because of the different directions of ion transport in the acid and alkali electrolytes, water was produced at the cathode with acid, and it was removed by the excess air flow. On the other hand, for an alkaline system, the water was produced at the anode, where there was no gas flow, so the only way to remove it would be to have it diffuse through the electrolyte to the cathode, where it could be removed by air.

Of the different acids, sulfuric was selected over phosphoric because, at the low operating temperatures (60-70°C) intended for the system, sulfuric acid had higher conductivity, and the oxygen electrodes used in tests performed slightly better in this acid.

Shell turned to the direct methanol-air fuel cell after deciding that their 5-kW demonstration hydrogen-air fuel cell system was too complex, operating on hydrogen produced from methanol and purified by a palladium-silver diffusion membrane. To minimize the corrosion of materials with an acidic electrolyte, the researchers chose to operate the cell at a low temperature, around 60°C, which would also minimize the methanol evaporation in the ambient pressure cell. At low temperatures, inexpensive plastics could be used as cell materials. For both anode and cathode, the catalyst could be applied to the electrode surface, which had a thin layer of gold as a conductor coated over a microporous polyvinyl chloride substrate (Williams et al., 1965). (The value of the gold-PVC substrate was $1.50/ft2.) Shell used platinum-ruthenium for the anode and platinum for the cathode (Andrew and Glazebrook, 1966). (This could be done on two sheets of plastic or on both sides of one sheet of plastic. The electrolyte would circulate on one or both sides of the fuel electrode.)

An eight-cell prototype was built in 1963 and produced 3.15 W at 1 A; after two years, it could still produce 2.85 W at 1 A, an indication of the durability of the materials. It used sulfuric acid electrolyte (6 N) with 1 M methanol mixed and circulated with electrolyte. A fuel cell stack with 40 cells was constructed and produced 300 W at 12 V and 60°C. In both fuel cell stacks, the circulating electro-lyte-methanol mixture emitted an ester-like odor, which upon analysis was found to contain formaldehyde and formic acid, intermediates that were expected, and also traces of acetic, propionic, butyric, and isobutyric acids, which had a source that had yet to be identified. These compounds had a poisoning effect on the anode catalyst, so a better catalyst was needed that could suppress the formation of these side reactions while also being effective for methanol oxidation.

The ESSO Research and Engineering Company developed direct methanol-air fuel cells with the goal of delivering a portable battery for military communications systems to the U.S. Army Electronics Laboratories (Tarmy and Ciprios, 1965). The work began in 1962, and by 1966 they had demonstrated a 60-W, 6-V unit that was self-sustaining. The cells used 3.7 M sulfuric acid electrolyte, and a control system added water to maintain the acid concentration. Methanol concentration was 0.75 M, which was found to be optimum, and it was controlled by a diffusion membrane that limited the diffusion to an electrode. The current was proportional to the concentration. The temperature range was 60 to 80°C, which could be regulated by the air flow rate that controlled the water removal. Ambient temperature air was supplied by a small blower and then humidified in a water economy unit that had a moisture-permeable membrane separating the countercurrent flows of fresh, ambient air and warm, wet exhaust air. The electrodes were supported by 52 mesh tantalum screens, 4 mil thick, on which 25 mg/cm2 of anode catalyst was pressed, and for the cathode, 9 mg/cm2 of platinum mixed with Teflon.

A demonstration unit produced 82 W at 6.0 V and 13.6 A, with the stack producing 99 W (7.0 V, 14.1 A), and the control systems consuming 15 W of the 17-W parasitic power loss. In continuous operation, the performance dropped to 60 W at 6.0 V.

At 60°C (140°F) for a single cell, terminal cell voltages of 0.50 V at 54 mA/cm2 (50 A/ft2) and 0.40 V at 108 mA/cm2 (100 A/ft2) were attained. In the 16-cell battery with larger electrodes, the performance was 50 mV less than the 0.46 and 0.31 V in the same size electrodes. In separate studies of a 16-cell module, Tarmy and Ciprios found that low methanol concentrations and high methanol conversion levels caused the electrodes to "starve," increasing polarization from 0.30 V at 4 vol% inlet methanol concentration to 0.44 V at 1 vol%. However, a lower methanol concentration could improve cathode performance because less methanol would migrate to the platinum cathode where it would be oxidized by the air, causing a mixed potential at the electrode and reducing the efficiency of the cell. Air flow was between two and ten times stoichiometric, variable in that range depending on the temperature control (water removal) needed.

Guide to Alternative Fuels

Guide to Alternative Fuels

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