hydroxylamine electrochemistry at polycrystalline platinum in acidic media: a voltammetric, dems and...

Journal ofElectroanalytical

Chemistry

Journal of Electroanalytical Chemistry 566 (2004) 53–62

www.elsevier.com/locate/jelechem

Hydroxylamine electrochemistry at polycrystalline platinum inacidic media: a voltammetric, DEMS and FTIR study

Victor Rosca *, Guillermo L. Beltramo, Marc T.M. Koper *

Laboratory of Inorganic Chemistry and Catalysis, Schuit Institute of Catalysis, Eindhoven University of Technology, 5600 MB Eindhoven,

The Netherlands

Received 7 November 2003; accepted 12 November 2003

Abstract

The electrochemical activity of hydroxylamine on polycrystalline Pt in acidic media has been characterized using cyclic vol-

tammetry, on-line differential electrochemical mass spectrometry (DEMS) and in situ FTIR. The electrochemistry of hydroxylamine

(HAM) is essentially controlled by other species that interact strongly with the electrode surface. Therefore, quite moderate current

densities, both in oxidation and reduction, are observed in a wide potential window between ca 0 and 1 V. The HAM electrore-

duction is a slow process and is masked by the Hupd. No formation of gaseous products was detected in this region; hence, ammonia

must be the main product of HAM reduction. The HAM electro-oxidation is strongly influenced by the adsorption of its products,

as well as their chemical transformations in solution. The key intermediate in HAM oxidation is NO, based both on voltammetric

and spectroscopic evidence. Nitric oxide forms an adsorbed layer, stable over a wide potential region between ca. 0.55 and 0.75 V.

At higher potentials NOads is oxidized to form (adsorbed) HNO2. At potentials above ca. 0.9 V the accumulation of HNO2 in

solution, accompanied by its partial oxidation to NO2, is postulated. N2O formation, observed in the potential region between ca.

0.5 and 1 V, has multiple sources. The most important source of N2O is a homogeneous reaction between HNO2 and HAM. The

Tafel slope analysis suggests the second electron transfer to be the rate-determining step in HAM oxidation to NOads. A tentative

mechanism for this reaction is proposed.

� 2003 Elsevier B.V. All rights reserved.

Keywords: Hydroxylamine; Polycrystalline platinum; Electrocatalysis; Voltammetry; DEMS; FTIR

1. Introduction

Hydroxylamine reactivity in an electrochemical en-

vironmental is important from a fundamental point of

view – in the context of the redox chemistry of inorganic

nitrogen compounds – as well as from a technological

point of view, mostly in relation to its industrial syn-

thesis. In industry, hydroxylamine (HAM), a key inter-

mediate in e-caprolactam production and a compound

with various applications in technology, is obtained bythe liquid-phase catalytic hydrogenation of nitric oxide

(NO) or nitrate at carbon-supported platinum or pal-

ladium catalysts [1,2]. The relevance of the electro-

* Corresponding authors. Tel.: +31-40-247-4916; fax: +31-40-245-

5054.

E-mail addresses: [email protected] (V. Rosca), M.T.M.Koper@

tue.nl (M.T.M. Koper).

0022-0728/$ - see front matter � 2003 Elsevier B.V. All rights reserved.

doi:10.1016/j.jelechem.2003.11.011

chemical studies on similar and model systems for

understanding the processes practiced in industry is wellaccepted [2].

This paper is a first step towards extending our pre-

vious mechanistic studies on NO electrochemistry on Pt

– so far oriented mainly toward the environmentally

important reduction to N2 or the complete reduction to

ammonia [3–5] – to investigating the factors controlling

selectivity toward HAM. However, the investigation of

these factors is hampered by a lack of experimental dataon HAM electrochemistry itself. This is particularly true

for HAM reduction on platinum electrodes in acidic

media [6], which are the conditions of interest from a

practical point of view.

The mechanisms of the electrochemical reactions of

HAM on Pt appear complex [6], and only few detailed

mechanistic studies have been reported so far. In a series

of studies, M€oller and Heckner [7,8] addressed the

mail to: [email protected]

54 V. Rosca et al. / Journal of Electroanalytical Chemistry 566 (2004) 53–62

reduction of HAM on bright Pt, using voltammetry and

the rotating disk electrode (RDE) technique. The ad-

sorption of HAM was shown to be the first step in HAM

reduction. The electroreduction on Pt proved difficult to

study, as it is strongly masked by the hydrogen ad-sorption and hydrogen evolution reaction (her). Thus,

M€oller and Heckner [7] interpreted their preliminary

results on HAM reduction at a Pt RDE in acidic media

(pH<3) by the hydrogen adsorption reaction:

NH3OHþads þ e� ! NH2OHþHads ð1Þ

In alkaline media HAM was much more active and re-

duction to ammonia was postulated, although with no

reference to the nature of the adsorbed intermediates. In

a later, more detailed, study, the authors argued for the

possibility of breaking the N–O bond in acidic media as

well. The rate-determining step was formulated as fol-

lows [8]:

NH3OHþads þ e� ! ðNH3Þads þ ðOH�Þads ð2Þ

The low reduction activity of HAM was explained by

the blocking effect of (NH3)ads.

M€oller and Heckner [9–12] also extensively studiedHAM electro-oxidation on bright Pt. Both NH2OH and

NH3OHþ were perceived as reactive species, depending

on pH. The HAM oxidation was described in terms of

successive dehydrogenation steps, paralleled by chemical

side reactions involving the reaction intermediates. In

acidic solution, the first two reaction steps of the

mechanism were formulated as follows:

NH3OHþads ! HNOHþ e� þ 2Hþ ð3Þ

HNOH ! NOH þ e� þHþ: ð4ÞBoth intermediates (which presumably may be consid-ered as adsorbed) were supposed to form N2, HNOH by

dimerization and disproportionation, and NOH by re-

action with HAM. Furthermore, by referring to the

homogeneous association–dissociation of the HNO/

H2N2O2 couple and the decomposition of H2N2O2 to

N2O (see [6] and references therein), the authors took

the experimentally observed N2O evolution as a proof

for NOH formation. It was suggested that nitrate is themain product of HAM oxidation in acidic media and

neutral media, while nitrite is perceived as a stable in-

termediate, particularly in alkaline media.

More recently, Karabinas et al. [13] presented a de-

tailed study of HAM anodic behavior on polycrystalline

Pt in neutral (buffered) and acid solutions by combining

voltammetry and on-line differential electrochemical

mass spectrometry (DEMS). In their mechanism, theauthors accepted the dehydrogenation scenario, result-

ing in NOH formation. The dimerization of NOH to

hyponitrous acid (H2N2O2) and its subsequent decom-

position to N2O and water were also suggested. At po-

tentials up to 0.7 V vs. RHE (pH 6.5, phosphate buffer)

NOH is partially oxidized to NO. As soon as Pt–OH

species become available at the surface, NOH was pro-

posed to be oxidized directly to HNO2, the latter being

considered to be a stable intermediate. The reaction of

HNO2 with HAM was held responsible for the forma-tion of N2O at potentials around 1 V. Above ca. 1.1 V

nitrite is oxidized to nitrate. Oxidation via Pt–OH spe-

cies was perceived as the main reaction pathway in

neutral solutions, while in acidic media the reaction can

partially proceed via NO2 formation.

The present paper describes a new detailed study on

HAM electrochemistry on Pt electrodes, oriented espe-

cially towards the effect of adsorption phenomena. Theelectrochemical data presented here are augmented by

results obtained with on-line DEMS – a technique al-

lowing detection of gaseous products. In order to iden-

tify adsorbates and (adsorbed) intermediates we have

also employed in situ FTIR spectroscopy. As a result,

we will put forward quite a detailed picture of HAM

transformations on polycrystalline Pt in acidic media, as

well as suggest some mechanistic interpretations.

2. Experimental

The H2SO4 and HClO4 working solutions were pre-

pared from their respective concentrated acids (‘‘Su-

prapur’’, Merck) and ultra-pure water (Millipore Milli-

Q system, 18.2 MX cm, 2 ppb total organic carbon). Theinfrared experiments aimed at detection of nitric oxide

and hydroxylamine were performed in H2SO4 solution

prepared with deuterium oxide (Merck, 99.8%). Hy-

droxylamine solutions were prepared by dissolving hy-

droxylammonium sulfate (‘‘Pro Analysi’’, Merck) in

ultra-pure water or deuterium oxide. Prior to each ex-

periment, all solutions were deoxygenated by purging

with argon.A platinum flag electrode was used as the working

electrode during the electrochemical measurements. For

the in situ FTIR measurements, a Pt disk electrode of 10

mm diameter was used. Prior to each experiment, the

working electrode was flame-annealed, cooled down in

an Ar atmosphere for ca. 20 s, quenched in ultra-pure

water and transferred to the cell.

Electrochemical measurements were performed in aconventional single-compartment three-electrode glass

cell, using a computer-controlled potentiostat (Auto-

Lab-PGSTAT20, Eco Chemie, Utrecht, The Nether-

lands). The cell and all the glassware were cleaned by

boiling in a 1:3 mixture of concentrated nitric and sul-

furic acids, followed by repeated boiling with ultra-pure

water. A platinum coiled wire served as the counter

electrode. In H2SO4 solutions a saturatedHgjHg2SO4jK2SO4 electrode, connected via a Luggin

capillary, was used as the reference. In HClO4 solutions

the reference electrode was an internal reversible hy-

Fig. 1. Cyclic voltammograms recorded at a polycrystalline Pt elec-

trode in the presence of hydroxylamine in solution. The numbers in the

legend indicate the upper potential limit of the corresponding cycle.

Experimental conditions: 0.5 M H2SO4; 5� 10�3 M NH2OH; poten-

tial sweep rate 5 mV s�1. Start potential 0.06 V.

V. Rosca et al. / Journal of Electroanalytical Chemistry 566 (2004) 53–62 55

drogen electrode (RHE). However, all potentials are

reported against the RHE.

On-line DEMS measurements were performed with a

computer-assisted Balzers Prisma QMS 200 mass spec-

trometer. The experimental setup was similar to thatdescribed elsewhere [14]. The working electrode, a

freshly platinized Pt gauze (with a roughness factor of

ca. 70–80), was laid on the hydrophobic Teflon mem-

brane (Schleicher and Sch€ull, pore size 0.02 lm), me-

chanically supported by a steel frit that separates the cell

from the vacuum chamber. The response time constant

was about 1 s. The mass signal for m=z 30 (NO) was

corrected for the ion fragmentation probability. Allmass intensities are given in arbitrary units.

The in situ Fourier transform infrared reflection ab-

sorption spectroscopy (FTIRRAS) measurements were

performed under external reflection conditions. The

spectrometer used was a Br€uker IFS113V, equipped

with a narrow-band MCT detector. The spectroelect-

rochemical cell [15] featured a prismatic CaF2 window

beveled at 60�. To acquire FTIRRAS data 500 inter-ferograms, at a resolution of 8 cm�1, were collected at

each potential. The reflectance spectra were calculated

as ðR� R0Þ=R0, where R and R0 are the reflectances at

the sample and reference potentials, respectively. Con-

sequently, the DR=R ratio gives negative bands for spe-

cies that are formed and positive bands for species that

are consumed at the sample potential, as compared to

the reference potential.

3. Results and discussion

3.1. Cyclic voltammetry

The voltammetric profile recorded at a Pt electrode at

low scan rate in the presence of HAM in solution isshown in Fig. 1. In order to avoid HAM oxidation, the

potential sweep experiments were always started in the

hydrogen underpotential deposition (Hupd) region at ca.

0.06 V. Within the limits of the Hupd region (approx.

0.06–0.4 V) HAM is apparently involved in reduction

process(es) only. These seem to be quite sluggish irre-

versible processes, as the increase of the potential sweep

rate up to moderate values (50–100 mV s�1) results in aprofile almost identical to the blank profile, though the

H adsorption peaks are somewhat broader and slightly

shifted towards negative potentials.

Remarkably, around the resting potential (ca. 0.42 V,

as determined from open circuit potential measure-

ments) both HAM reduction and oxidation may take

place simultaneously, which is in agreement with the fact

that HAM may disproportionate in the presence of a Ptcatalyst [1]. Hence, there is no well-defined boundary

between the HAM electroreduction and electro-oxida-

tion. A gradual increase of the upper potential limit

gives rise to the oxidation feature EO4, followed by a

region (0.55–0.75 V) of low positive current. Above ca.

0.75 V another oxidation feature emerges (EO5). In the

negative-going part of the cyclic voltammograms, pre-

sented in Fig. 1, an important increase in height and a

slight change in position of the reduction peak ER2, as

well as the development of the feature ER1 are observed,

while peak ER3 does not change. Similar experiments inperchloric acid give a qualitatively identical picture

(Fig. 2). However, slightly higher positive current den-

sities are observed in the potential region between 0.4

and 0.9 V, most probably due to the absence of any

significant anion adsorption, resulting also in lower

potentials at which the surface oxidation occurs. The

qualitative effect of anion co-adsorption was well de-

fined in all experiments performed in perchloric acid, inspite of introducing a small amount of sulfate by adding

hydroxylammonium sulfate.

These voltammetric features can be explained in

terms of HAM oxidation to NO (anodic peak EO4),

resulting in the formation of an NO ad-layer, and its

subsequent reductive stripping in the negative-going

sweep (cathodic features ER1, ER2, ER3). The following

considerations support this conclusion. First, the posi-tion of the three reduction features is very close to those

observed in the case of the reductive stripping of an NO

ad-layer under the same experimental conditions [3].

Second, the Tafel slope analysis of the peaks ER2 and

ER3, when plotted as peak position ðEpÞ vs. logarithm of

the scan rate (log(m)) [16,17], gives slopes of ca. 50 and

19 mV dec�1 (Fig. 3), respectively, values close to those

observed for the corresponding features in the case of

Table 1

The results of the least-square linear regression analysis of the peak

current vs. the potential sweep rate dependence for the features (peaks)

presented in Fig. 1

Peak H2SO4 HClO4

105 (intercept�SDa)

(A s V�1)

105 (intercept�SDa)

(A s V�1)

ER2 )5.60� 0.88 )5.94� 0.16

ER3 )8.43� 0.18 )11.84� 0.02

EO4 3.79� 0.52 3.26� 0.40

Experimental conditions: as in Fig. 1. 5–10 experimental points

were used in regression analysis. The square of the correlation coeffi-

cient ðR2Þ was higher than 0.999 in all cases.a Standard deviation of the intercept.

Fig. 2. Cyclic voltammograms recorded at a polycrystalline Pt elec-

trode in the presence of hydroxylamine in solution. The numbers in the

legend indicate the upper potential limit of the corresponding cycle.

Experimental conditions: 0.5 M HClO4; 5� 10�3 M NH2OH; poten-

tial sweep rate 5 mV s�1. Start potential 0.06 V.

Fig. 3. Tafel slope analysis of the reduction peaks ER2 and ER3 (see

Fig. 1). Experimental conditions: 0.5 M H2SO4; 5� 10�3 M NH2OH.

The CV profiles analyzed were obtained according to the following

potential program: (i) after a potential of 0.7 V was applied for 10 s, (ii)

the potential was swept in the negative direction.


NO ad-layer stripping (54 and 20 mV dec�1) [3]. Im-portantly, the ER3 peak should be considered as a vol-

tammetric feature corresponding to the hydrogen

adsorption with some contribution of the co-adsorbed

HAM and the adsorbed NO reduction (if any NO was

generated at higher potentials). However, the value of

the Tafel slope, though of no mechanistic value, is likely

to be determined by the presence of the adsorbed NO.

Once formed, the NO ad-layer essentially blocks further

HAM oxidation in the potential region between 0.55

and 0.75 V. Nevertheless, from the effect of the upper

potential limit on the reduction peak ER2 and reductionfeature ER1, it can be concluded that in this region slow

oxidation of HAM and NO ad-layer formation are still

taking place.

Further evidence for adsorbed NO as the main

(stable) intermediate of HAM oxidation is provided by

transfer experiments. The working electrode was im-

mersed in a deoxygenated electrolyte solution (sulfate)

containing 10 mM of HAM for 1 min. Next, the elec-trode was thoroughly washed with ultra pure water and

immediately transferred to another electrochemical cell

containing no HAM. Then, the electrode potential was

swept in the negative-going direction. The resulted

voltammetric profile strongly resembles that observed

in the case of the NO ad-layer stripping [3]. Further-

more, a single potential sweep was enough to recover

the blank voltammogram of clean platinum. The NOcoverage, determined from the charge analysis, ranged

between 0.2 and 0.3. In conclusion, adsorbed NO is the

only adsorbed intermediate detected in the transfer

experiments.

From a linear regression analysis of the peak height

ðIpÞ vs. potential sweep rate ðmÞ dependence (Table 1) it

can be concluded that the processes corresponding to

features ER2, ER3 and EO4 are all surface confined, bothin sulfuric and perchloric acid. For the features ER2 and

ER3 the intercepts (extrapolated to zero scan rate) are

slightly different from zero in both media, as the NO

reductive stripping is paralleled by the H adsorption and

slow HAM reduction. Similarly, the small non-zero in-

tercept in case of the oxidation feature EO4, also ob-

served by Karabinas et al. [13] in neutral (buffered)

media, may point to the existence of a slow parallel(continuous) oxidation process.

Fig. 4 shows voltammetric profiles recorded at dif-

ferent HAM concentrations. As can be seen from the

data for the lowest concentration, the positive-going

Fig. 4. The effect of the hydroxylamine concentration on the voltam-

metric profile recorded at a polycrystalline Pt electrode. Experimental

conditions: as in Fig. 1. The hydroxylamine concentration as indicated

by the numbers in the legend.

Fig. 5. The Tafel slope analysis for the oxidation peak EO4 in sulfuric

acid. Experimental conditions: as in Fig. 1.

Fig. 6. The Tafel slope analysis for the oxidation peak EO4 in per-

chloric acid. Experimental conditions: as in Fig. 2.


scan, started at 0.06 V, shows a gradual increase of the

negative current within the Hupd region, with a maxi-

mum at ca. 0.3 V. This could imply that the reduction of

HAM may be simply blocked by the Hupd, as is also the

case for the nitrate reduction on Pt [18].

The Tafel slope analysis of the oxidation feature EO4

gives values of 41 and 38 mV dec�1 in sulfuric andperchloric acid, respectively, thus suggesting the second

electron transfer to be rate-limiting in the oxidation

process. Interestingly, in a narrow window of sweep

rates, the Tafel slope linearity breaks down (Figs. 5 and

6), both in sulfuric and perchloric acid. At higher sweep

rates the slope apparently re-establishes its value as that

observed at low sweep rates. This effect is more pro-

nounced for data obtained in perchloric acid (Fig. 6).Again, this points towards to the occurrence of a sur-

face-confined process in parallel with a (slow) continu-

ous oxidation process. We will discuss a possible

candidate for this latter process in the next two sections.

3.2. On-line DEMS measurements

Differential electrochemical mass spectrometry is atechnique enabling detection and, under certain condi-

tions, measuring the formation rate of gaseous as well as

volatile products as a function of the applied electrode

potential [14,19]. The ion current is recorded simulta-

neously with the cyclic voltammetry run. The resulting

curves are usually referred to as mass spectrometric

cyclic voltammograms (MSCVs). The technique is par-

ticularly useful in the case of electrochemical transfor-mations of inorganic nitrogen-containing compounds,

as the formation of various gaseous products is possible

[3,4,13,18,20].

The cyclic voltammograms recorded at the DEMS

working electrode, along with the corresponding

MSCVs for selected m=z ratios, recorded both in sulfuric

and perchloric acid, are presented in Figs. 7 and 8, re-spectively. It should be mentioned that evolution of the

m=z 28 signal (N2), also shown in Figs. 7 and 8, is dif-

ficult to interpret in terms of in situ electrolytic forma-

tion of N2, as the background signal was too high for

this mass. In sulfuric acid there is a wide potential

window between ca. 0 and 0.8 V where hardly any

gaseous products could be detected in the positive going

sweep (see Fig. 7). First, this result indicates that nosignificant formation of gaseous products takes place in

Fig. 8. Mass signal intensity (in arbitrary units) vs. potential and

current density vs. potential curves recorded simultaneously at Pt-

black gauze electrode. Experimental conditions: 0.5 M HClO4;

5� 10�2 M NH2OH; potential sweep rate 2 mV s�1. Start potential

0 V.

Fig. 7. Mass signal intensity (in arbitrary units) vs. potential and

current density vs. potential curves recorded simultaneously at a

Pt-black gauze electrode. Experimental conditions: 0.5 M H2SO4;

5� 10�2 M NH2OH; potential sweep rate 2 mV s�1. Start po-

tential 0 V.


the Hupd region, leaving ammonia as the only reduction

product. Additionally, this result is in good agreement

with voltammetric data, which points to the formation

of an NO ad-layer and its blocking effect on further

HAM oxidation in potential region between ca. 0.6 and

0.8 V. At potentials above 0.8 V, there is a parallel in-


crease of the electrical current density and the formation

of NO (m=z 30), N2O (m=z 44) and NO2 (m=z 46).In order to facilitate further discussion of the DEMS

results, it will be useful to summarize briefly the mech-

anisms of NO oxidation on Pt. The final product of theadsorbed NO oxidation on Pt in acidic media is nitrate,

while nitrite, or nitrous acid, is perceived as an impor-

tant intermediate [21]. However, adsorbed NO proved

to be quite resistant to complete oxidation, the process

being accompanied by the oxidation of the Pt surface

[22]. From the Tafel slope analysis, an EC mechanism

for the oxidation of adsorbed NO on Pt was proposed.

The following electrochemical equilibrium is assumed:

NOads þH2O¡ðHNO2Þads þHþ þ e� ð5ÞThe subsequent rate-determining chemical step is at

present unclear. Kinetic analysis, based on results from

rotating disk experiments, indicates that continuous NO

oxidation proceeds via HNO2 to nitrate as the finalproduct [23].

Having established the formation of HNO2, N2O

formation at potentials higher than ca. 0.8 V (Fig. 7, m=z44) can be explained by a reaction of HNO2 with HAM

present in solution, as also argued by Karabinas et al.

[13] in the case of HAM electro-oxidation in neutral

(buffered) media. Curiously, formation of some N2O is

also observed well below 0.8 V in the negative-goingpart of MSCVs with a maximum at ca. 0.4 V. As the

simultaneously recorded CV shows low current densities

in the region under consideration, this could still be a

homogeneous reaction between residual HNO2 and

HAM, though the maximum at ca. 0.4 V suggests the

existence of an additional source of N2O formation,

possibly electrochemical. Interestingly, based on DEMS

measurements, Nishimura et al. [24] argued for the di-rect electrolytic HNO2 reduction to N2O in a potential

region between 0.2 and 0.8 V vs. RHE in 0.5 M H2SO4.

In a more recent FTIR study, Bae et al. [25] have con-

firmed this conclusion.

In strong acids (see [6] and references therein), NO2 is

perceived as the main product of HNO2 oxidation and

the primary intermediate of HNO2 oxidation to nitrate

in moderately acidic (pH 0.5–3) and neutral solution[26]. This could account for the onset of the mass signal

corresponding to m=z 46 (NO2) at potentials just above

0.8 V (Fig. 7). The MSCV at m=z 30 is non-zero even

after correction for the ion fragmentation. This fact

points to accumulation of some NO in solution. Possible

sources of NO could be homogeneous disproportion-

ation of HNO2 (see [6] and references therein) or, per-

haps, partial desorption of the NO ad-layer.The MSCVs recorded in perchloric acid give a picture

qualitatively similar to that in sulfuric acid, though

higher electrical current densities as well as mass signals

for the 46, 28 and 30 m=z values are observed (Fig. 8).

This is probably due to the lower potentials at which

oxygen adsorption on Pt begins in perchloric acid (see

the blank CV in Fig. 2). The NO accumulation in so-

lution starts at ca. 0.7 V, ca. 100 mV lower compared to

sulfuric acid. Additionally, two peaks in NO formation

(m=z 30) can be seen, suggesting multiple sources of NO.Interestingly, formation of some N2O in perchloric acid

is observed at potentials as low as ca. 0.5 V (see the inset

in the plot for m=z 44, Fig. 8) in the positive-going

sweep, compared to ca. 0.8 V in sulfuric acid. A possible

explanation for this result is a slow catalytic reaction

between HAM and adsorbed NO, as suggested by van

de Moesdijk [1] from catalytic studies on carbon-sup-

ported Pt, although admittedly the effect of the catalystpotential and electrolyte composition are not known.

M€oller and Heckner [7] and later Karabinas et al. [13]

have suggested a dimerization of NOH on the surface

and the subsequent disproportionation of the dimer as

the possible source of N2O. However, our own results

have shown that the reduction of adsorbed NO does not

yield N2O unless there is NO in solution [3].

When presenting CV data, we have argued for am-monia to be the main, if not the only, product of HAM

reduction in the Hupd region. However, as suggested in

early studies on electrochemical reduction of HAM on

mercury [27,28] and later argued in a detailed mecha-

nistic study by Heyrowsky and Vavricka [29], HAM

may be involved in enhancing the hydrogen reduction.

In order to check the occurrence of this effect on Pt, we

have followed the influence of the HAM concentrationon the m=z 2 (H2) mass signal. The applied potential was

0.06 V, i.e. the negative potential limit of Hupd region. A

gradual increase of the HAM concentration indeed

brought about a (statistically) significant gradual in-

crease of the m=z 2 mass signal. A leveling of the mass

signal was observed at ca. 50 mM HAM. The mass

current reached was ca. 3% higher than that in the ab-

sence of NH2OH in solution. Hence, the effect may bestatistically significant but certainly not catalytically

significant.

3.3. In situ FTIR measurements

The consumption of HAM as well as the accumula-

tion of the intermediates and products of its oxidation in

the thin layer and on the platinum surface can bemonitored as a function of the applied potential by in

situ FTIRRAS, once their characteristic spectroscopic

features are well separated from the others. The band

displayed at ca. 1460 cm�1 [30] was assigned to the –

NOH bending ðdÞ mode in NH3OHþ and used for

HAM detection. This assignment is confirmed by values

calculated with the Gaussian 98 software package [31],

predicting a value of 1507 cm�1 for the –NOH bendingðdÞ mode in H3NOHþ and 1486 cm�1 in D3NOHþ.Other possible vibrational modes are well separated

from the 1460–1500 cm�1 frequency range (see Table 2).

Table 2

Vibration frequences corresponding to different vibrational modes in hydroxylamine and ammonia in UHV, as calculated with the Gaussian 98

software package or as cited in the literature

Species Vibration frequency (cm�1) Ref.

ðdÞ-N–H2 ðdÞ-N–O–H ðmÞ-N–O–

H3NOHþ 1598–1686 1507 1196 This work

D3NOHþ 1205–1219 1486 991

D3NODþ 1205–1219 1159 988

NHþ4 1680 – – [31]

NDþ4 1215 – –


Note that –(H–N–H) ðdÞ mode exhibits an isotopic shift

of 400 cm�1 compared to –(D–N–D) (Table 2). An

isotopic effect is also expected for the –N–O–H ðdÞmode, when H is replaced by D (Table 2). However, this

effect is not expected, as in our experiments we have

used fully hydrogenated HAM, and the H atom in –

NOH group is not exchangeable. Bulk N2O gives anasymmetric stretch ðm1Þ band, centered at 2231 cm�1

[32]. In principle, detection of NO, both in adsorbed and

dissolved forms, is possible, though the assignment of

the vibrational modes of NO adsorbed on the Pt surface

must be carried out with great care. In all IR experi-

ments to be described below, the upper potential limit

was 0.9 V, in order to avoid massive HAM depletion in

the thin layer. In all cases the reference spectrum was thespectrum collected at 0.1 V.

Fig. 9 shows the IR spectra for selected frequency

ranges and electrode potentials. From Fig. 9(A) and (B)

consumption of HAM and formation of N2O in the thin

layer can be deduced. The band frequencies of adsorbed

Fig. 9. Potential difference FTIR spectra collected at a polycrystalline Pt elec

0.1 V. Experimental conditions: 0.1 M H2SO4 + 1.5� 10�2 M NH2OH; in H

NO are situated in a wide frequency range between 1450

and 1800 cm�1, depending on both coverage and applied

potential (see for example [33] and references therein).

As shown in Fig. 9, spectrum (C), in this frequency

window a series of negative-going spectroscopic features

is observed. Some of these, in particular the feature(s) in

the 1550–1650 cm�1 frequency range, can be tentativelyassigned to the vibrational modes of multifold coordi-

nated NO [34]. Features in the 1550–1650 cm�1 range as

well as a bipolar band centered at ca. 1760 cm�1 (as

compared to the negative-going band at ca. 1720 cm�1

in Fig. 9(C)) were observed by Gootzen et al. [35] for

NO ad-layers on polycrystalline platinum.

The integrated area of the bands corresponding to

different vibrational modes provides a measure of theconsumption or formation of that species in the thin layer.

By plotting the integrated area vs. potential, trends in

HAM consumption and N2O formation can be followed.

From such a plot (Fig. 10), it can be deduced that HAM is

consumed over the whole potential region examined.

trode at different applied potentials (as indicated). Reference potential

2O (B) and in D2O (A and C); p-polarized light.

Fig. 10. Trends in hydroxylamine consumption and N2O formation as

a function of the applied potential. See text for details. Experimental

conditions: as in Fig. 9.


Apparently,HAMconsumption slows down at potentials

where the EO4 oxidation feature occurs, thus confirming

the surface blocking effect of adsorbed NO, as also de-

duced from the voltammetric profile (Fig. 1). From thesame plot it can be seen that N2O formation starts at ca.

0.45 V, close to the potential at which N2O formation is

observed in perchloric acid by DEMS (Fig. 8). This result

allows us to conclude that N2O formation starts at ca.

0.45 V in both sulfuric and perchloric acid, although the

extent of its formation is controlled by the nature of the

anion. This is a strong indication for the N2O formation

being the (slow) continuous oxidation process parallel toNOads formation, as from Figs. 5 and 6 this process is

expected to begin around 0.5 V.

3.4. Mechanistic implications for HAM oxidation

The results of the Tafel slope analysis suggest the

second electron transfer to be the rate-determining step

in HAM oxidation to surface-bonded NO. The follow-ing simplified reaction scheme is tentatively proposed:

H3NOHþ¡H3NOHþ

ads ð6Þ

H3NOHþads¡H2NOads þ 2Hþ þ e� ð7Þ

H2NOads ! HNOads þHþ þ e� rds ð8Þ

HNOads ! NOads þHþ þ e� ð9ÞThe HNO and H2NO intermediates are suggested on

the basis of their lower energy in vacuum compared to the

NOH and HNOH species, as deduced from theoretical

calculations [3,5]. However, it has to be admitted that the

stability of surface-bonded species is not known. It was

also suggested [5] that in the back reaction to reaction (9)

the proton is transferred directly from the solution. Thisidea is also supported by the position of the oxidation

feature EO4, which occurs at a potential where the H

coverage is negligible, if not zero. We have assumed here

that the protonated form of HAM is the active species,

although it is commonly assumed that both NH2OH and

NH3OHþ are reactive species, depending on pH (see [6]

and references therein). As discerned from the DEMS as

well as FTIR results, starting at ca. 0.5 V, the process is

paralleled by the (slow) formation ofN2O,most probablythrough a reaction between adsorbedNOandHAMfrom

the solution. At potentials above ca. 0.8 V, adsorbed NO

is subjected to further oxidation.

It is informative to compare the reaction scheme

proposed for HAM oxidation to NOads and the mech-

anism proposed for reduction of NOads [3,5]. The NOads

reduction is perceived as a stepwise hydrogenation

process, in which the formation of HNO and H2NOintermediates is assumed:

NOads þHþ þ e�¡HNOads equilibrium ð10Þ

HNOads þHþ þ e� ! H2NOads rds ð11Þ

H2NOads þ 4Hþ þ 3e� ! NHþ4 þH2O: fast ð12Þ

Ammonia was shown to be themain product of reduction

of the adsorbed NO on platinum. Importantly, the pro-

tons in reactions (10) and (11) seem to be transferred di-

rectly from a solution hydronium species [5]. Comparingreactions (8) and (11), it is noticed that these reactions are

the forward and back reactions of the same transforma-

tion and, at the same time, rate-determining steps in the

corresponding reaction schemes. This suggests that the

energetic barrier for theHNO$H2NO transformation is

the highest barrier in the whole NO$NH2OH chain of

transformations, whatever the direction.

3.5. Mechanistic implications for HAM reduction

The HAM electroreduction on platinum is strongly

masked by the Hupd and the her at higher overpotentials.

Most importantly, the experimental data presented

above strongly suggest the absence of any formation of

gaseous products in the Hupd region, thus pointing to

ammonia as the main, if not the only, product of HAMcathodic activity:

NH3OHþads þ 2e� þ 2Hþ ! NHþ

4 þH2O: ð13ÞAt high hydrogen coverage the HAM reduction is in-

hibited, as the HAM adsorption sites are blocked by theHupd. The nature and adsorption geometry of the ad-

sorbate corresponding to the nitrogen oxidation state in

HAM is currently not known. Finally, a catalytic effect

of HAM toward the hydrogen reduction process on

platinum, as claimed for Hg electrodes, does not appear

significant.

4. Conclusions

In this paper we have reported on the electrochemical

activity of hydroxylamine on polycrystalline Pt in acidic


media, as discerned from cyclic voltammetry, on-line

DEMS and in situ FTIRRAS.

Since HAM appears to adsorb relatively weakly on

platinum, HAM electrochemistry seems to be controlled

by other species that interact strongly with the electrodesurface, e.g. hydrogen or (bi)sulfate anion. This can be

seen even better from studies on single-crystal Pt elec-

trodes [36]. Therefore, quite moderate current densities,

both in oxidation and reduction, are observed in a wide

potential window between ca 0 and 1 V.

The HAM electroreduction is a slow process and is

strongly masked by the Hupd. At the same time, HAM

reduction is possible in the Hupd region only. We did notobserve formation of gaseous products in this region;

hence, ammonia must be the main product of HAM

reduction.

The HAM electro-oxidation is strongly influenced by

the adsorption of its products, as well as their chemical

transformations in solution. The key intermediate in

HAM oxidation is NO. It forms an adsorbed layer,

virtually stable over a wide potential region between ca.0.55 and 0.75 V. It acts as a poison for further HAM

oxidation, in a fashion very similar to the way adsorbed

CO poisons methanol oxidation. At higher potentials,

where oxygen adsorption starts, NOads is oxidized to

form (adsorbed) HNO2. At potentials as high as ca. 0.9

V the accumulation of HNO2 in solution, accompanied

by its partial oxidation to NO2, are postulated. N2O

formation, observed in the potential region between ca.0.5 and 1.0 V, probably has multiple sources. The most

important source of N2O is a homogeneous reaction

between HNO2 and HAM [37], since a dramatic increase

in N2O formation is observed at potentials where NO is

oxidized to HNO2.

The Tafel slope analysis suggests the second electron

transfer to be the rate-determining step in HAM oxi-

dation to NOads, the latter being perceived as the stableintermediate in the total oxidation. Our results are lar-

gely in agreement with the mechanistic picture given by

Karabinas et al. [13]. Firstly, we have also observed N2O

formation starting at ca. 0.5 V. Furthermore, the

mechanism of HAM oxidation to NO, proposed in this

paper, assumes formation of two intermediates, corre-

sponding to the HAM dehydrogenation steps. The main

difference between our mechanism and that suggestedpreviously lies in the identification of the key adsorbed

intermediate. Whereas Karabinas et al. have suggested

NOH, our results seem to point to NO as the stable

intermediate in HAM oxidation, based on voltammetric

as well as spectroscopic evidence.

Acknowledgements

This work was supported by a grant from the Neth-

erlands Organization for Scientific Research (NWO).

Technical assistance by Jos van Wolput in the IR mea-

surements is gratefully acknowledged.

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hydroxylamine electrochemistry at polycrystalline platinum in acidic media: a voltammetric, dems and...

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