Fuel Cell Systems

20 Anodic Oxidation of Derivatives of Methane, Ethane, and Propane in Aqueous Electrolytes

Downloaded by CHINESE UNIV OF HONG KONG on June 18, 2016 | http://pubs.acs.org Publication Date: January 1, 1969 | doi: 10.1021/ba-1965-0047.ch020

I.

Galvanostatic Investigations

H. BINDER, A. KÖHLING, H. KRUPP, K. RICHTER, and G . SANDSTEDE

Battelle-Institut e.V., Frankfurt(Main), Germany

Anodic potential/current density plots of a Raney platinum electrode were measured galvanostatically with alcohols, aldehydes, ketones, and carboxylic acids up to 800 mv. (vs. hydrogen electrode in the same solution) at 25° and 80°C, the electrolyte being 5N potassium hydroxide and 5N sulfuric acid. Acetic acid and propionic acid cannot be oxidized in potassium hydroxide solution but are active in sulfuric acid. With methanol, ethanol, glycol, and glycerol in 5N potassium hydroxide at 200 ma./sq.cm. potentials from 300 mv. to 350 mv. observed, while in 5N sulfuric acid the potentials range from 500 to 650 mv. under the same conditions.

Qrganic substances may be used as electrochemical fuels provided that they can be oxidized both easily and rather completely. It is therefore interesting to determine the reaction rate and the reaction mechanism of suitable compounds. Derivatives of hydrocarbons are not only potential intermediates in the anodic oxidation of hydrocarbons, but they themselves might be used as fuels. A number of papers (1, 2, 4-8) mainly dealt with the mechanism of anodic methanol oxidation. The reactivities of additional organic compounds were examined particularly by Pa vela (5) and Schlatter (7). In a previous paper (3) we reported on the application of electrodes containing Raney-type transition metals as catalysts i n the electrochemi269 Young and Linden; Fuel Cell Systems Advances in Chemistry; American Chemical Society: Washington, DC, 1969.

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FUEL CELL SYSTEMS

cal oxidation of hydrogen and methanol i n aqueous electrolytes. Raney platinum proved to be a very active catalyst. W e have now studied the anodic behavior of alcohols, aldehydes, ketones, and carboxylic acids with 1 to 3 carbon atoms i n potassium hydroxide solution and sulfuric acid at different temperatures. G a l vanostatic potential/current density plots were measured, and a coulometric-potentiostatic method was employed for quantitative analysis.


Experimental Conditions Reactants and Electrodes. for our experiments: Methanol Formaldehyde Formic acid

W e selected the following 12 substances

Ethanol Acetic acid Glycol Oxalic acid

1-Propanol 2JPropanol Propionic acid Acetone Glycerol

These reactants were dissolved i n the electrolyte. In general, the individual substances were used i n concentrations of 2 moles/liter, the solvent being aqueous solutions of 5N sulfuric acid or 5 N potassium hydroxide. In the carboxylic acids, equivalent quantities of potassium hydroxide were added when the solvent was basic. The electrodes were prepared i n the form of porous disks by compressing a mixture of gold powder (99.96%, particle size < 60 micron) forming the electrode skeleton, the powdered Raney platinum aluminum alloy and sodium chloride as an auxiliary substance to obtain macropores. Subsequently, the aluminum of the Raney alloy was dissolved by treating the electrode first with dilute and then with concentrated potassium hydroxide solution, at a final temperature of 80°C. The Raney platinumaluminum alloy was PtALi. In some runs with methanol and formic acid, we used electrodes with Raney palladium; i n these cases the Raney alloy was P d A U . The Raney alloys were prepared by heating an intimate mixture of the catalyst powder and aluminum powder in the form of compressed disks. Then the alloys were powdered and sieved and the 25 to 40 micron size fraction of the powdered alloy was used i n an amount of 35% by volume i n the test electrode. The electrode was suspended by a platinum wire in the electrolyte. The current densities mentioned here always refer to the projected surface area of the electrode. Measuring Setup. Figure 1 shows the half-cell arrangement for the galvanostatic measurements. The test electrode is the Raney platinum anode. A platinum wire served as counter electrode. As reference electrode we used an "autogenous" hydrogen electrode, as we would like

Young and Linden; Fuel Cell Systems Advances in Chemistry; American Chemical Society: Washington, DC, 1969.

20.

BINDER ET AL

Oxidation of Paraffin Derivatives.

I.

271


to call this type of hydrogen electrode. In our experiments this electrode is a Raney platinum electrode suspended by a platinum wire i n a glass tube, which has a capillary tube attached to it. O n this Raney platinum electrode, hydrogen is being evolved at a cathodic current density of about 1 ma./sq.cm. A t this current density the overvoltage for hydrogen evolution is negligibly small on this electrode, so that it may be used as an accurate hydrogen reference electrode. For every solution and temperature employed the potential of the electrode is assumed to be zero.

autogenous Hy^ \ reference electrode]

_J>^

capillary

\ \Pt counter electrode test electrode

Figure 1. Half-cell arrangement Hence, the potentials of the anode depicted i n the following figures relate to the potential of a hydrogen electrode i n the same solution. The ohmic drop between capillary and electrode is not accounted for. A l l values were taken under quasi steady-state conditions; they remained constant during the measuring period of sometimes more than 24 hours. The plots were taken first at decreasing and then at increasing current densities. This implies that a certain proportion of the reaction products is present i n the electrolyte. Measurements in 5N KOH Solution Methane Derivatives. Figure 2 shows the results of the measurements with methanol, formaldehyde, and formic acid i n 5N potassium hydroxide at 25°C. The larger rate of oxidation is obtained for methanol and formaldehyde, while formic acid furnished inferior results. The peculiar shape of the curve for formaldehyde at small current densities is caused b y formaldehyde dehydrogenating on the Raney

Young and Linden; Fuel Cell Systems Advances in Chemistry; American Chemical Society: Washington, DC, 1969.

272

FUEL CELL SYSTEMS

platinum catalyst. In this case, therefore, we are dealing with the oxi dation of hydrogen, rather than the conversion of formaldehyde. Only at larger current densities, where the dehydrogenation rate is no longer large enough, is the formaldehyde converted directly. WOO I I Η COOK >!

/


(+;

3

[HCHO

(χ)

500

C

C

\CH OH

οQ.

0)

2 M Reactor t 5 Ν KOH 25°C (H ) 2

50

WO

150

200

current density -ma I cm *

Figure 2.

Potential/current density plots of a Raney platinum electrode with methane derivatives in 5N KOH at 25°C.

WOO

2M Reactan 5 Ν KOH 80° C

1

500-

Η COOK \(CH OH f

Q.

Fuel Cell Systems

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