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Chapter 1. Introduction 8 For CO2 electrolysis is about 1.47 V. For H2O electrolysis, the H2O must be supplied to the cell electrode as steam. This lowers the from about 1.48 V (water electrolysis) to 1.3 V (steam electrolysis) since part of the energy has already been supplied in the phase transition H2O(l) H2O(g). However, unless steam is already available, the water must be vaporized to steam by the system upstream of the cell, and if this is to be done by electrical heating then the cell–vaporizer system can isothermally be operated at the 1.48 V with a heat exchanger included to transfer the heat being produced by the cell (the cell at 1.48 V produces heat because it is operated about 0.18 V above the ―true‖ ) to the vaporizer. A heat exchanger is needed anyway to transfer heat from the outlet gasses of the cell to the inlet gasses in order to conserve heat. Because heat exchangers are not 100% efficient and the insulation around the cell stack is also imperfect, isothermal operation of the system requires operating somewhat above to accommodate heat losses. If 10% of the heat produced is lost to the surroundings, the cell must be operated at 10% higher voltage. See also section 3.2.2 with Figure 3-4 for more information about the thermodynamics including the enthalpy and Gibbs free energy values, and section 3.3 for more discussion about cell operation. 1.1.2. Kinetics A typical current density versus voltage (i-V) polarization curve of a solid oxide cell at 850 °C, with 50% H2 + 50% H2Oiii supplied to the negative-electrode and oxygen or air supplied to the positive-electrode, is depicted in Figure 1-3 with the above described thermodynamic potentials labeled. The cell voltage is scanned downwards (resulting in fuel cell operation, producing current) and upwards (resulting in electrolysis operation, drawing current) from the OCV. The voltage difference between the applied potential and OCV is referred to by various terms – applying polarization, bias, or overpotential. These three terms are generally used interchangeably in this thesis. One can see in this hypothetical cell that for a given change in voltage, slightly more current is produced by fuel cell operation than is consumed in electrolysis. This means that, because the current density is directly proportional to the electrochemical reaction ratesiv , the reaction rate is faster for fuel cell operation (in this case H2 oxidation) than electrolysis operation (in this case H2O reduction). This is observed in real cells as well, as will be seen for the cell studied in Chapter 4. The net reaction rates can be described by the area-specific resistance (ASR) which is the change in cell voltage divided by the change in iii 50% H2 + 50% H2O is one useful gas composition for characterizing the cells operated reversibly in fuel cell and electrolysis mode. In real cell operation, the gasses supplied would be closer to 100% H2 for fuel cell operation and 100% H2O for electrolysis operation (although some electrode materials, including the traditional negative electrode material, are unstable in 100% H2O so some H2 is also supplied, which can be recirculated H2 that was produced by the electrolysis reaction). iv As defined by Faraday’s law, assuming that a negligible amount of current is due to parasitic reactions, which is known to be true for the majority of fuel cells/electrolyzers including solid oxide cellsPDF Image | Electrolysis of CO2 and H2O
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