Electrochemical Oxidation of Ethanol
J. Phys. Chem. B, Vol. 101, No. 2, 1997 257
platinum electrodes. The decrease in acetic acid production with
an increase in surface step density coincides with trends
observed in infrared spectroscopy experiments.9,10 During a
potential sweep from -0.25 to +0.6 V (vs. SCE), vibrational
bands for acetic acid appeared at less positive potentials on Pt-
(111) than on other platinum surfaces. In these spectroscopic
studies, acetic acid formation was blocked at Pt(110), Pt(100),
and polycrystalline platinum electrodes at potentials below about
0.4 V, and the two electron oxidation product acetaldehyde
appeared dominant.9,10 This inhibition in the acetic acid
producing pathway at stepped platinum electrodes (Pt(100)
electrodes have been shown to restructure in aqueous solution,
forming nanoscopic domains with high step site coverages26)
coincides with the decrease in acetic acid formation at Pt(557)
and Pt(335) electrodes indicated in Figure 4a.
The absolute amounts of acetic acid detected by each
technique can also be compared. Over the potential range
between 0.25 and 0.35 V, infrared spectroscopy determinations
indicate that 2-3 nmol of acetic acid are generated during a 2
mV/s sweep (50 s measurement period).9,10 Figure 4a shows
that 1.2 nmol of acetic acid are produced during a 60 s
measurement period at 0.3 V. Agreement between the two
methods is within a factor of 2, which approaches the uncertainty
in the electromagnetic enhancement (ref 9, p 323) in the thin
layer cavity.9,10 Currently, the protocol used in chromatographic
determination of electrochemically generated acetic acid is
limited by uncertainty in the solution volume of the drop that
protects the single crystal electrode during transfer to the
electrolyte. With more accurate measurement of drop volumes,
the chromatographic technique can provide a means for calibra-
tion in quantitative infrared spectroelectrochemical experiments.
The conditions employed in the present study did not allow
for detection of acetaldehyde and carbon dioxide, which are
also major reaction products. Strategies for determination of
aldehydes, through reaction with phenylhydrazine compounds,30
are under investigation. In addition, it should be noted that
acetaldehyde can oxidize to acetic acid in the presence of
atmospheric oxygen.31-33 Our studies of this side reaction have
indicated that it is slow and does not significantly affect the
acetic acid determinations under the present experimental
conditions.
The chromatography results are also consistent with product
distributions from ethanol oxidation determined with the use
of in-situ mass spectrometry. Acetaldehyde was determined to
be the main two carbon oxidation product in reactions at
polycrystalline platinum,27,28 platinum black,1 and platinum-
ruthenium1 electrodes, which have a high density of low
coordination sites. Acetic acid has not been detected as a
product of ethanol oxidation in in-situ mass spectrometry studies,
which thus far appear restricted to polycrystalline and Pt(110)
electrodes.1,3,4,27,28
The surface structure sensitivity of ethanol oxidation is due
in part to facile C-C bond cleavage at step sites.11 The faster
C-C bond breaking kinetics lead to larger CO2 yields9-11 but
at the expense of increased poisoning by adsorbed CO9-11 and
possibly other partial oxidation products.3,4,12,13,28 The decrease
in acetic acid production with increased surface step density is
likely a consequence of (i) a shift in reactions toward the CO2
producing pathway and (ii) increased surface poisoning that
blocks sites for water activation (eq 2).9,10 The latter can account
for the preferential formation of acetaldehyde on partially
blocked stepped surfaces.9,10
Conclusions
The ethanol oxidation pathway leading to the formation of
acetic acid is markedly affected by the electrode atomic-level
surface structure. The quantity of acetic acid produced during
positive potential sweeps is greatest across the range 0.2-0.6
V and decreases with increasing surface step density. Studies
show that soluble products of reactions at single crystal
electrodes can be isolated and quantified with the use of ion
chromatography. In the determination of soluble species, the
chromatographic technique can overcome limitations of in-situ
infrared spectroscopy and provide a means for calibration in
infrared experiments.
The chromatographic experiments also provide a check on
the quantitative capabilities of the thin layer infrared spectros-
copy technique. Infrared spectroelectrochemistry has many
vitrues, especially in the detection of adsorbates;18,29 however,
the thin layer cell geometry can limit accuracy in the quantifica-
tion of soluble species. Electric field inhomogeneity in the thin
layer cavity,14-16,18,19 product leakage,10 reactant depletion,10,17
and nonuniform cavity thickness17 can affect the intensity of
vibrational bands and lead to errors in quantitative measure-
ments.16 Also, in the study of ethanol oxidation, the carbonyl
stretching bands for acetaldehyde and acetic acid overlap, and
independent determinations require extending the infrared
Acknowledgment. We are grateful for the assistance of Dr.
G. M. Swain (Utah State University) during the fabrication of
single crystal electrodes, Dr. N. Gu¨ven and V. Polyak with X-ray
backscattering measurements, and Drs. C. E. Evans (University
of Michigan), P. K. Dasgupta, and A. Sjo¨gren with preliminary
ion chromatography experiments. We also appreciate the
assistance of Mr. J. Serafini with acetaldehyde stability studies.
Financial support for this work was provided by the U.S. Office
of Naval Research.
spectral window below 1000 cm-1 9,10
Ion chromatography
.
allows the soluble products to be determined more specifically
(vide infra) and quantitatively.
References and Notes
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(2) Beden, B.; Leger, J.-M.; Lamy, C. In Modern Aspects of Electro-
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New York, 1992; Vol. 22, p 97.
(3) Iwasita, T.; Pastor, E. Electrochim. Acta 1994, 39, 531.
(4) Schmiemann, U.; Muller, U.; Baltruschat, H. Electrochim. Acta
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(5) Cases, F.; Lopez-Atalaya, M.; Vazquez, J. L.; Aldaz, A.; Clavilier,
J. J. Electroanal. Chem. 1990, 278, 433.
(6) Morin, M.-C.; Lamy, C.; Leger, J.-M.; Vasquez, J.-L.; Aldaz, A.
J. Electroanal. Chem. 1990, 283, 287.
(7) Cases, F.; Vazquez, J.-L.; Perez, J. M.; Aldaz, A. J. Electroanal.
Chem. 1991, 310, 403.
(8) Cases, F.; Morallon, E.; Vazquez, J.-L.; Perez, J. M.; Aldaz, A. J.
Electroanal. Chem. 1993, 350, 267.
(9) Leung, L. W. H.; Chang, S. C.; Weaver, M. J. J. Electroanal. Chem.
1989, 266, 317.
Figure 4a can be compared to the plots of acetic acid quantity
vs potential, generated from in-situ infrared spectra.9,10 The
curve for Pt(111) in Figure 4a corresponds to the traces for acetic
acid in Figure 9 of ref 10 and Figure 3 of ref 9. The curve in
Figure 4a is proportional to the deriVatiVe of the plots generated
from infrared spectra. Since infrared spectra were collected
during a slow potential sweep, products accumulate in the thin
layer cavity through the duration of the scan. In contrast,
samples for chromatography were obtained following a single
potential step and are not affected by the buildup of products
generated at intermediate potentials. Data from Figure 4a and
infrared spectroscopy show excellent correspondence. Both
indicate the rate of acetic acid generation maximizes between
0.3 and 0.4 V and then quickly falls off for potentials e 0.4 V.