Potential Dependence of the Yield of Carbon Dioxide from Electrocatalytic Oxidation of Methanol on Platinum(100)

Article abstract of DOI:10.1021/jp9629248

The yield of complete oxidation product (CO2) from electrocatalysis of methanol on a Pt(100) electrode in perchloric acid electrolyte at room temperature was determined by combined measurements of chronoamperometry and linear sweep voltammetry.The fractional yield of CO2 was zero over a potential range of 0.35-0.45 V_RHE and increased monotonically to unity over the potential range of 0.5-0.65 V.The results may be explained by a simple parallel mechanism, in which methanol oxidizes directly to CO2, or by a complex serial mechanism, in which the overpotential for CO oxidation is reduced by solution phase methanol.

Full text of DOI:10.1021/jp9629248

© Copyright 1997 by the American Chemical Society
VOLUME 101, NUMBER 19, MAY 8, 1997
LETTERS
Potential Dependence of the Yield of Carbon Dioxide from Electrocatalytic Oxidation of
Methanol on Platinum(100)†
T. D. Jarvi, S. Sriramulu, and E. M. Stuve*
UniVersity of Washington, Department of Chemical Engineering, Box 351750, Seattle, Washington 98195-1750
ReceiVed: September 23, 1996; In Final Form: February 21, 1997X
The yield of complete oxidation product (CO2) from electrocatalysis of methanol on a Pt(100) electrode in
perchloric acid electrolyte at room temperature was determined by combined measurements of chronoamper-
ometry and linear sweep voltammetry. The fractional yield of CO2 was zero over a potential range of 0.35-
0.45 VRHE and increased monotonically to unity over the potential range of 0.5-0.65 V. The results may be
explained by a simple parallel mechanism, in which methanol oxidizes directly to CO2, or by a complex
serial mechanism, in which the overpotential for CO oxidation is reduced by solution phase methanol.
Introduction
to be CO, forms as a serial reaction intermediate, reaction 1
Development of a low-temperature, direct methanol fuel cell
requires substantial improvements in the kinetics of methanol
electrooxidation. One major problem with this reaction is
poisoning of the surface by accumulation of CO1-6 and other
carbonaceous species such as HCO or COH.7-10 On typical
catalysts like carbon-supported platinum or bimetallic platinum/
ruthenium, oxidative removal of adsorbed carbon monoxide
requires potentials much greater than the desired potential of a
technical anode. The higher activity of Pt/Ru electrocatalysts
in comparison with pure Pt is due partly to reduced poisoning
levels brought about by ruthenium-enhanced oxidation of carbon
monoxide.11-13 Yet even for the most active bimetallic
catalysts, anode poisoning remains a problem.
+H2O, –2H+
–4H+
–4e–
r1s
CH3OH
CO
CO2
(1)
–2e–
r2s
or as a side product in a parallel reaction mechanism, eq 2.
–4H+
CO
–4e–
r1p
CH3OH
(2)
r2p
+H2O, –6H+
–6e–
CO2
Despite considerable research on methanol electrooxidation,
a definitive description of the reaction mechanism remains
elusive. The main issue is whether the poison, here assumed
Equivalently, one may ask whether methanol can oxidize
completely to CO2 at potentials where CO does not oxidize.
From a technical perspective, the nature of further research in
this area depends upon the answer to this question. If reaction
proceeds by the serial route, then further research (if any) must
attempt to destabilize adsorbed CO. If, instead, the parallel path
is operative, then further research should focus on improving
the kinetics and selectivity to CO2, reaction r2p.
† Presented at the 212th ACS National Meeting, Orlando, FL, 1996; Paper
COLL-51.
* Address correspondence to this author. E-mail: Stuve@
u.washington.edu. Fax: 206-543-3778.
X Abstract published in AdVance ACS Abstracts, April 15, 1997.
S1089-5647(96)02924-0 CCC: $14.00 © 1997 American Chemical Society

3650 J. Phys. Chem. B, Vol. 101, No. 19, 1997
Letters
No clear consensus exists as to which mechanism holds.
Early results1,14 were interpreted in favor of the parallel
mechanism, yet the surfaces used in those studies were not
sufficiently well characterized in terms of cleanliness, number
of active sites, surface structure, or accumulation of poison.
More recent work with single-crystal substrates and time
dependent reaction measurements (chronocoulometry, chrono-
amperometry) does address these concerns. Through compara-
tive studies of carbon monoxide and methanol oxidation on Pt/
Ru electrodes, Gasteiger et al.11,12 found in favor of the serial
path, whereby they attributed the Ru-induced enhancement for
methanol reaction to increased CO oxidation kinetics. Vielstich
and Xia15 similarly found in favor of the serial path as their
DEMS (differential electrochemical mass spectrometry) mea-
surements on platinum electrodes showed methanol reacting
first to CO (or other carbonaceous species) without production
of CO2 for potentials below 0.45 V. Conversely, Herrero et
al.16,17 found in favor of the parallel path. Their argument rests
on turnover numbers that exceed unity on Pt(100) at potentials
where adsorbed carbon monoxide is stable. The conclusion
here is that some methanol must react completely to CO2 (r2p)
without formation of CO (r1p); otherwise, the surface would
saturate with CO, and the turnover number could not exceed
unity.
Figure 1. Cyclic voltammogram (sweep rate of 50 mV/s) of flame-
annealed Pt(100), cooled in Ar.
In this Letter we present electrochemical measurements of
the fraction of methanol that oxidizes completely to CO2 as a
function of the potential of a Pt(100) electrode. Reaction studies
were performed at constant potential for a fixed period of time
(60 s). Accumulation of CO during reaction was measured
by linear stripping voltammetry performed immediately after
the reaction period. A simple charge balance allows for
calculation of the CO2 yield from measured charges, and the
resulting yield data are discussed in terms of their mechanistic
significance.
Experimental Details
The experiments were performed in a dual electrochemical
cell apparatus operating in an anaerobic environment. The two
electrochemical cells were identical, and each consisted of a
Teflon electrolyte cup of 0.25 cm3 volume, platinum foil counter
electrode of 0.4 cm2 surface area, and palladium R-hydride
reference electrode. The electrochemical cells were similar to
those previously described,18 and complete details will appear
in a later publication.
Figure 2. Schematic representation of the multielectrolyte experiment.
The upper traces show the current measured for the potential wave
forms shown in the lower traces. The arrow denotes electrode transfer
from cell A to cell B. Respective concentrations (M) of HClO4 and
CH3OH were (a) 0.1, 0.1 and (b) 0.1, 0.
Single-crystal platinum beads were prepared by melting a 0.5
mm diameter platinum wire (ESPI, 99.99%) in a hydrogen-
oxygen flame to form a bead approximately 1.5 mm in
diameter.19 The beads were etched in aqua regia at 60 °C for
30 min to reveal a faceting characteristic of the (100) and (111)
crystal faces. The (100) face was aligned by visual inspection
and polished with diamond paste to 1 µm smoothness. After
polishing, the samples were annealed at 1100 °C in a hydrogen-
air flame for 30 min.20 Before each electrochemical measure-
ment, the electrode was heated to about 1000 °C for several
seconds and quenched in argon-saturated, ultrapure water. The
sample was protected by a droplet of ultrapure water during
the subsequent transfer to the electrochemical cell. A cyclic
voltammogram for samples prepared in this fashion is shown
in Figure 1, from which we conclude that the surface is free of
contamination and ordered along the (100) plane. Figure 1
compares well with previously reported voltammograms for
flame-annealed Pt(100).21,22 The relative heights of the peaks
at 0.285 and 0.35 V depend on the exact cooling procedure.21,22
In particular, the lower potential peak increases in height at the
expense of the higher potential peak if sample transfer from
the flame to the quenching water takes too long.
The CO2 yield measurements were performed as follows.
Cell B of the two-cell system contained blank electrolyte (0.1
M HClO4), and cell A contained methanol-bearing electrolyte
(0.1 M CH3OH in 0.1 M HClO4) or blank electrolyte for rinsing
the electrode. After a quick anneal the clean electrode was
immersed into cell B for recording of its cyclic voltammogram.
The electrode was emersed at 0.075 V and immersed into cell
A at the same potential. After 2 s the potential was stepped to
the desired reaction potential Er for 60 s while the resulting
current transient was recorded as depicted in Figure 2a. The
charge passed during the current transient is the reaction charge
qr. After the reaction period the electrode potential was returned
to 0.075 V. Traces of methanol-bearing electrolyte were
removed by rinsing the electrode, still in cell A, with blank
electrolyte under potential control. The electrode was then
transferred to cell B for measurements of adsorbed, partial
oxidation products by linear sweep voltammetry as depicted in
Figure 2b. The charge measured under the CO oxidation peak
at 0.72 V is the stripping charge qs. The yield of CO2 was
determined from qr and qs as described below.

Letters
J. Phys. Chem. B, Vol. 101, No. 19, 1997 3651
Figure 3. Reaction charge (a) qr (0, 9) and the stripping charge (b)
qs (O, b) as a function of reaction potential Er. Open symbols represent
data for cumulative turnovers of less than one and solid symbols for
cumulative turnovers greater than one. The size of data points reflects
(1 standard deviation in the qr measurement. The lines are guides to
the eye.
Figure 4. Fractional yield yc of CO2 as a function of reaction potential.
The line is a guide to the eye. Details of the yield calculation are
given in the text. Typical standard deviations are shown for some data
points.
as shown in eq 2, though the final form of the equation does
not depend on the choice of mechanism.24 The number of moles
of methanol that react during the reaction period is denoted Nr,
the Faraday constant F, and the electrode surface area A. The
complete oxidation product is CO2 and we assume that the
partial oxidation product is CO, so that six electrons are required
for complete oxidation and four electrons for partial oxidation.
The reaction charge qr can be expressed as
Electrolyte solutions were prepared from perchloric acid
(GFS, doubly distilled), methanol (Fisher, spectrophotometric
grade), and ultrapure water (Barnstead Nanopure, 18 MΩ cm
resistivity). Both reference electrodes were prepared identically
by charging at -150 µA for 15 min daily. Potential control
and voltammetry were performed with an EG&G 263 poten-
tiostat. Current transients were recorded by a Tektronix TDS
520 digital oscilloscope. Measured reaction transient currents
were at least 2 orders of magnitude less than diffusion-limited
currents estimated by the Cottrell equation.23 For all electro-
chemical measurements only the polished face of the electrode
contacted the meniscus of the electrolyte. A drop of blank
electrolyte was maintained on the electrode during cell transfers
to protect the surface from contamination. All measurements
were performed at room temperature, and electrode potentials
are reported with reference to a reversible hydrogen electrode
in 0.1 M HClO4.
F
A
qr ) Nr(6yc + 4yp)
(3)
If the CO formed by r1p in eq 2 is oxidized according to r2s in
eq 1, the stripping charge required is
F
A
qs ) 2 Nryp
(4)
Since the yields of complete and partial oxidation products are
related by
Results and Discussion
yc + yp ) 1
(5)
The reaction charge qr and the stripping charge qs are shown
as a function of the reaction potential Er in Figure 3. The
reaction charge is approximately constant below 0.45 V but
increases at 0.5 V and beyond. The qs values scatter about
a line at approximately 235 µC/cm2 below 0.65 V, while at
higher potential they drop off. The drop in the stripping charge
above 0.65 V is attributed to an increase in the rate of CO
oxidation (r2s) during the reaction period τ at this potential
(compare with Figure 2b). Data corresponding to cumulative
turnovers less than or greater than unity are depicted with
open and filled symbols, respectively; the surface supports
complete turnover for Er g 0.5 V. Cumulative turnovers were
calculated on the basis of a six-electron transfer process, a
reaction site of one surface atom, and the density of topmost
atoms of an ideal Pt(100) surface (1.28 × 1015 cm-2); a reaction
charge of 1.23 mC cm-2 corresponds to a cumulative turnover
of one.
we may rearrange eqs 3-5 to obtain an expression for the
fractional yield of complete oxidation product
qr - 2qs
qr + qs
yc )
(6)
The fractional yield as determined by eq 6 is shown in Figure
4 as a function of reaction potential. The results show three
distinct potential regions corresponding to different product
distributions. At low potential (below 0.50 V), partial oxidation
products form exclusively, in agreement with DEMS experi-
ments.15 Reaction at intermediate potentials (0.5-0.65 V)
produces both CO and CO2. Finally, beyond the apparent onset
of CO oxidation (0.65 V), CO2 is the dominant reaction product.
Figure 4 indicates CO2 production occurs at potentials of 0.5-
0.65 V, while CO is predicted to be surface stable at these
potentials by either Figure 2b or Figure 3. Furthermore, Figure
4 predicts CO2 as the majority product over this entire potential
range.
A simple analysis leads from the data reported in Figure 3 to
an expression for the fractional yield yc of complete oxidation
product (CO2). Here fractional yields of complete yc or partial
yp oxidation products are defined as the moles of product formed
per mole of methanol reacted during the reaction period τ. To
develop the necessary equation, we assume a parallel mechanism
These results may be interpreted according to either the
parallel or serial mechanism. The parallel mechanism provides
the most direct explanation as the increase in yc above 0.5 V

3652 J. Phys. Chem. B, Vol. 101, No. 19, 1997
Letters
can be interpreted as an increase in r2p relative to r1p.
Alternatively, the serial mechanism also explains the data if one
assumes that CO oxidation through r2s occurs at appreciable
rates over the potential range of 0.5-0.65 V. Such an
assumption appears to contradict the voltammetric measurement
in Figure 2b, which shows CO to be surface stable below 0.65
V. However, the reaction charge was measured in a methanol-
containing electrolyte, whereas the stripping charge was mea-
sured in the absence of methanol. Solution phase methanol
could provide the chemical potential necessary to reduce the
overpotential of reaction r2s relative to that in methanol free
electrolyte.
To examine the serial mechanism hypothesis further, note
that the tailing edge of the reaction transient in Figure 2a implies
the existence of a near steady state. The serial mechanism
predicts a steady state when r1s ) r2s. At the end of the reaction
period t ) τ, the rate of CO oxidation r2s(τ) must therefore be
finite according to the serial mechanism and measured turnover
numbers exceeding unity. To test this, we performed one
reaction pulse on the electrode at 0.55 V for 60 s in methanol-
bearing electrolyte (cell A) and then transferred the electrode
to blank electrolyte (cell B) and performed a second reaction
pulse, again to 0.55 V for 60 s. If r2s were finite, then the current
transient at the beginning of the second pulse should be about
one-third that at the end of the first pulse. However, we detected
no measurable current transient (other than double-layer charg-
ing), and the reaction charge for the second pulse was below
the limit of detection. A subsequent stripping voltammogram
revealed that the CO coverage was the same as in the case where
only a single reaction pulse was performed. This result shows
that CO does not oxidize appreciably at 0.55 V so that the only
manner in which the serial mechanism could hold in this
potential range is in the case that solution phase methanol alters
the reaction mechanism in some other fashion than that shown
in eq 1. In other words, the mechanism for CO oxidation below
0.65 V (if any) is different than that above 0.65 V.
interpreted in favor of a serial mechanism, while our detection
of complete oxidation above 0.5 V along with turnover numbers
in excess of unity agrees with the previous results and
interpretation in favor of the parallel mechanism.16 Additional
preliminary results suggest that the parallel mechanism may
operate even at very low potentials.25 The simple parallel
mechanism, as represented in eq 2, indeed provides the most
straightforward interpretation of our results. A simple serial
mechanism, as shown in eq 1, cannot describe our results.
Instead, a more complex serial mechanism is required in which
solution phase methanol reduces the overpotential for CO
oxidation. We are continuing our experiments in this area in
order to determine which of the simple parallel or complex serial
mechanisms best represents electrooxidation of methanol on
platinum.
Acknowledgment. This work benefited from stimulating
discussions with H. Baltruschat, C. Campbell, T. Engel, N.
Markovic, and A. Wieckowski. We gratefully acknowledge
financial support from the National Science Foundation (CTS-
9502971).
References and Notes
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The assumption that the only partial oxidation product formed
is CO appears in eqs 3 and 4, where the numbers of electrons
required for the formation and oxidation of the partial oxidation
product are 4 and 2, respectively. Evidence of other stable
reaction products exists, namely HCO or COH species.9 The
presence of any finite quantity of HCO or COH species as partial
oxidation products leads to a wider potential window over which
yc is nonzero. This can be seen by considering the analog of
eq 6 for the case of either COH or HCO as the partial oxidation
product, for which the formation and oxidation steps both require
three electrons
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qr - qs
qr + qs
yc )
(7)
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For the same measured qr and qs values, the calculated yc values
are greater with this assumption. Thus, the yc values shown in
Figure 4 represent the lower bound to the yield of complete
oxidation products.
These results indicate that methanol electrooxidation on Pt-
(100) produces CO2 at potentials where CO appears to be surface
stable (0.5-0.65 V). Our inability to detect complete oxidation
below 0.5 V is consistent with the DEMS results,15 which were