**3.1 Ag/AgCl reference electrode**

Base cyclic voltammograms for a high-surface-area Pt-coated Au disk in 0.5 mol dm-3 KOH solution as a function of temperature are shown in Fig 2. Analogous to literature results, three characteristic zones corresponding to hydrogen electrochemistry, double layer and oxygen electrochemistry are observed. The peaks associated with hydrogen adsorption/desorption in the hydrogen electrochemistry zone over −1.15 to around −0.60 V are negatively shifted as the temperature is increased, accompanied by obvious decrease in the peak currents at higher temperature. The decrease in these peak currents and the double layer currents observed in a narrow potential range over -0.60 V to -0.50 V suggest that the electrochemical surface area of the electrodes is decreased at higher temperature. This could be caused by the accelerated adsorption of electrolyte components and/or impurities onto the electrode surface, resulting in the blocking of some electrode surfaces. In the oxygen electrochemistry zone over −0.5 to approximately 0.2 V, the potential of the cathode peak corresponding to the reduction of surface oxides is slightly shifted as the temperature is increased from 20° to 150°C. All these facts are indicative of the possibility of using the Ag/AgCl electrode as the internal reference electrode in aqueous KOH solution at elevated temperature.

The standard potential of the Ag/AgCl electrode in aqueous solution containing chloride has been measured in a wide temperature range over 0 to approximately 300 °C, and its temperature dependence has been analyzed by several groups (Öijerholm et al., 2009; Greeley et al., 1960). Based on the standard potential of the Ag/AgCl electrode of around 0.23 V vs. standard hydrogen electrode (SHE) and the reversible hydrogen potential in 0.5 mol dm-3 KOH solution (−0.81 V vs. *SHE*) at 20 °C, the value of the onset potential for hydrogen evolution would be around −1.04 V vs. Ag/AgCl if the hydrogen evolution

wt% Pd/C) were used for the preparation of gas diffusion electrodes with carbon-cloth as

During the measurements of base voltammograms and methanol electrooxidation, highpurity nitrogen was introduced into the electrochemical cell to inhibit vaporization of the liquid phase at elevated temperature, and gas-phase pressure was set at 300 psi unless otherwise stated. Background voltammograms were measured in 0.5 mol dm-3 KOH at 50 mV s-1. All methanol oxidation experiments were carried out in 0.5 mol dm-3 KOH + 0.5 mol dm-3 methanol. Steady-state voltammograms for the methanol oxidation were recorded at a scan rate of 10 mV s-1. The measurements of the chronoamperograms were performed by stepping potential from -1.1 V where no methanol oxidation occurs to a given value where methanol is oxidized. For the purpose of comparison, the oxidation of hydrogen was investigated under similar conditions by introducing high-purity hydrogen into the electrochemical cell to equilibrate with the solution of 0.5 mol dm-3 KOH solution at 200 psi. All oxygen reduction experiments were carried out in 0.5 mol dm-3 KOH solution equilibrated with 300 psi high purity O2. To investigate the electrochemical oxidation of CO, high purity CO was introduced into the electrochemical cell with the pressure set at 300 psi.

Base cyclic voltammograms for a high-surface-area Pt-coated Au disk in 0.5 mol dm-3 KOH solution as a function of temperature are shown in Fig 2. Analogous to literature results, three characteristic zones corresponding to hydrogen electrochemistry, double layer and oxygen electrochemistry are observed. The peaks associated with hydrogen adsorption/desorption in the hydrogen electrochemistry zone over −1.15 to around −0.60 V are negatively shifted as the temperature is increased, accompanied by obvious decrease in the peak currents at higher temperature. The decrease in these peak currents and the double layer currents observed in a narrow potential range over -0.60 V to -0.50 V suggest that the electrochemical surface area of the electrodes is decreased at higher temperature. This could be caused by the accelerated adsorption of electrolyte components and/or impurities onto the electrode surface, resulting in the blocking of some electrode surfaces. In the oxygen electrochemistry zone over −0.5 to approximately 0.2 V, the potential of the cathode peak corresponding to the reduction of surface oxides is slightly shifted as the temperature is increased from 20° to 150°C. All these facts are indicative of the possibility of using the Ag/AgCl electrode as the internal reference electrode in aqueous KOH solution at elevated

The standard potential of the Ag/AgCl electrode in aqueous solution containing chloride has been measured in a wide temperature range over 0 to approximately 300 °C, and its temperature dependence has been analyzed by several groups (Öijerholm et al., 2009; Greeley et al., 1960). Based on the standard potential of the Ag/AgCl electrode of around 0.23 V vs. standard hydrogen electrode (SHE) and the reversible hydrogen potential in 0.5 mol dm-3 KOH solution (−0.81 V vs. *SHE*) at 20 °C, the value of the onset potential for hydrogen evolution would be around −1.04 V vs. Ag/AgCl if the hydrogen evolution

the diffusion layer. All gases used were of research grade.

**2.3 Electrochemical measurements** 

**3. Results and Discussion 3.1 Ag/AgCl reference electrode** 

temperature.

overpotential on Pt is neglected. This value is very close to our measured value (−1.1 V). This agreement between the calculated value and measured value suggests that the Ag/AgCl electrode satisfactorily functions as the internal reference electrode in hydroxide solution. This is also supported by the characters of the hydrogen and oxygen electrochemistry for Pt electrode at elevated temperature.

Fig. 2. Cyclic voltammograms for a high-surface Pt-coated Au disk electrode of 0.5 mm diameter in 0.5 mol dm-3 KOH at a scan rate of 50 mV s-1 as a function of reaction temperature.

To further evaluate the suitability of using the Ag/AgCl electrode as the internal reference in the wide temperature range of our interest, the cyclic voltammograms for a high-surfacearea Pt-coated Au disk in 0.5 mol dm-3 KOH solution equilibrated with 200 psi H2 are measured and shown in Fig. 3. The positive-going scans are characteristic of fast-rising currents at lower overpotentials, followed by slow-rising currents and limiting currents at higher overpotentials. These characteristics are similar to literature results for H2 oxidation at Pt electrodes in aqueous base solution (Bao & Macdonald, 2007; Schmidt et al., 2002). Increasing reaction temperature clearly increases the limiting currents. These changes are mainly caused by the concentration changes of dissolved H2 as a function of the temperature. The potential value at zero current (*EH2,I=0*) measured in Fig. 3 is negatively shifted by around 40 mV as the temperature is increased from 20 to 150 °C.

The reversible potential of hydrogen reaction depends upon the reaction temperature and the pH value of the reaction medium as follows:

$$E\_{H\_2}^{o} = \frac{-2.303RT}{F}pH\tag{2}$$

where *T* is in K. Substituting pH=13.69 (corresponding to 0.5 mol dm-3 OH¯) and the values of *R* and *F* into Equation 2 results in its simplification as follows:

$$E\_{H\_2}^{\circ} = -2.7 \times 10^{-3} T \tag{3}$$

Investigations of Intermediate-Temperature Alkaline Methanol


difference between Ag/AgCl electrode and *RHE*.

**3.2 Methanol electrooxidation** 

intermediate-temperature range.


**EAg/AgCl-ERHE**

potentials can be corrected to the RHE scale according to Equation 4.

**ERHE**


**Potential / V**



Fuel Cell Electrocatalysis Using a Pressurized Electrochemical Cell 167

0 50 100 150 200

**C**

**Temperature / o**

Fig. 4. Temperature dependence of calculated *RHE* potential and measured potential

In Equation 4, the first term is much smaller than the second term being a constant. This supports that the temperature dependence of the Ag/AgCl electrode is similar to that of the RHE. Therefore, the Ag/AgCl electrode can be approximated as an equivalent of the RHE in preliminary investigations. The variations of measured electrode potentials with temperature are signature of the changes in the reaction kinetics. For the kinetic analysis, all

Cyclic voltammograms for a high-surface-area Pt-coated Au disk electrode in 0.5 mol dm-3 KOH solution containing 0.5 mol dm-3 methanol as a function of temperature (Fig. 5) clearly show that the onset potentials of substantial methanol electrooxidation are negatively shifted by increasing temperature. Because this prominent negative potential shift with increasing temperature is not caused by the potential shift of the reference electrode, it is reasonably believed that the electrooxidation of methanol is substantially accelerated. To assess the kinetics of methanol electrooxidation, we have compared the onset potentials for methanol oxidation and H2 oxidation (Fig. 3) under similar conditions. The value of their onset potential difference is significantly decreased with increasing temperature. At 150 °C, a typical value is approximately 60 mV. It is well accepted that the hydrogen oxidation is highly facile on Pt, this value indicates that highly facile electrooxidation of methanol can be achieved on single-element Pt electrocatalyst in aqueous alkaline solution in the

Fig. 3. Temperature dependence of cyclic voltammograms for a high-surface Pt-coated Au disk electrode of 0.5 mm diameter in 0.5 mol dm-3 KOH equilibrated with 200 psi H2 at a scan rate of 10 mV s-1.

According to Equation 3, the reversible hydrogen potential (*RHE*) should be negatively shifted with increasing temperature. At 20°C, the value is −0.81 V, and it is expected to be −1.14 V at 150°C. This means that a negative reversible potential shift should be around 0.33 V as the reaction temperature is increased from 20° to 150°C if the potential of the Ag/AgCl reference electrode is a constant independent upon the temperature. Because the value of *EH2,I=0* is a signature of the reversible hydrogen potential, its shift with increasing temperature should be theoretically similar to that of the *RHE* if the Ag/AgCl reference electrode potential remains constant. However, Fig. 3 demonstrates that increasing temperature from 20° to 150°C produces a negative shift of only around 40 mV, much lower than 0.33 V. Therefore, the small *EH2,I=0* shift strongly indicates that the potential of the Ag/AgCl reference electrode would have a temperature dependence similar to that of the *RHE*. This consistence provides a big convenience to investigate the kinetics of fuel cell reactions in a wide temperature range using an internal Ag/AgCl reference electrode since the Ag/AgCl reference electrode could be used as an equivalent of the reversible hydrogen electrode.

The potential difference between the Ag/AgCl electrode and the *RHE* is experimentally measured in a H2 atmosphere. Fig. 4 shows the temperature dependence of the potential difference. It is clearly seen that the difference is less dependent upon the reaction temperature in comparison to the individual *RHE*. A potential difference of around 80 mV as the temperature is increased from 20 to 150 °C. The temperature dependence of their potential difference could be linearly approximated as follows:

$$
\Delta E = E\_{Ag/AgCl} - E\_{RHE} = -4.4 \times 10^{-4} T - 0.980 \tag{4}
$$


 60 <sup>o</sup> C 80 <sup>o</sup> C 105 <sup>o</sup> C 130 <sup>o</sup> C 150 <sup>o</sup> C

**Potential / V vs Ag/AgCl**

Fig. 3. Temperature dependence of cyclic voltammograms for a high-surface Pt-coated Au disk electrode of 0.5 mm diameter in 0.5 mol dm-3 KOH equilibrated with 200 psi H2 at a

According to Equation 3, the reversible hydrogen potential (*RHE*) should be negatively shifted with increasing temperature. At 20°C, the value is −0.81 V, and it is expected to be −1.14 V at 150°C. This means that a negative reversible potential shift should be around 0.33 V as the reaction temperature is increased from 20° to 150°C if the potential of the Ag/AgCl reference electrode is a constant independent upon the temperature. Because the value of *EH2,I=0* is a signature of the reversible hydrogen potential, its shift with increasing temperature should be theoretically similar to that of the *RHE* if the Ag/AgCl reference electrode potential remains constant. However, Fig. 3 demonstrates that increasing temperature from 20° to 150°C produces a negative shift of only around 40 mV, much lower than 0.33 V. Therefore, the small *EH2,I=0* shift strongly indicates that the potential of the Ag/AgCl reference electrode would have a temperature dependence similar to that of the *RHE*. This consistence provides a big convenience to investigate the kinetics of fuel cell reactions in a wide temperature range using an internal Ag/AgCl reference electrode since the Ag/AgCl reference electrode could be used as an equivalent of the reversible hydrogen

The potential difference between the Ag/AgCl electrode and the *RHE* is experimentally measured in a H2 atmosphere. Fig. 4 shows the temperature dependence of the potential difference. It is clearly seen that the difference is less dependent upon the reaction temperature in comparison to the individual *RHE*. A potential difference of around 80 mV as the temperature is increased from 20 to 150 °C. The temperature dependence of their

> 4 / 4.4 10 0.980 *EE E Ag AgCl RHE <sup>T</sup>* (4)

potential difference could be linearly approximated as follows:




**Current / A**

scan rate of 10 mV s-1.

electrode.

**EH2 ,I=0**

0.0000

0.0001

Fig. 4. Temperature dependence of calculated *RHE* potential and measured potential difference between Ag/AgCl electrode and *RHE*.

In Equation 4, the first term is much smaller than the second term being a constant. This supports that the temperature dependence of the Ag/AgCl electrode is similar to that of the RHE. Therefore, the Ag/AgCl electrode can be approximated as an equivalent of the RHE in preliminary investigations. The variations of measured electrode potentials with temperature are signature of the changes in the reaction kinetics. For the kinetic analysis, all potentials can be corrected to the RHE scale according to Equation 4.
