Oxidation of Hydroxylamine on Gold Electrodes in Aqueous

Article pubs.acs.org/JPCC

Oxidation of Hydroxylamine on Gold Electrodes in Aqueous Electrolytes: Rotating Ring-Disk and In Situ Infrared Reflection Absorption Spectroscopy Studies Adriel Jebin Jacob Jebaraj, Doe Kumsa, and Daniel A. Scherson* Ernest B. Yeager Center for Electrochemical Sciences and Department of Chemistry, Case Western Reserve University, Cleveland, Ohio 44106-7078, United States S Supporting Information *

ABSTRACT: The oxidation of hydroxylamine on Au electrodes in aqueous phosphate buffer solutions (pH 7) was examined using electrochemical and in situ infrared reflection absorption spectroscopy techniques. Polarization curves recorded with a rotating Au disk electrode showed that the onset of NH2OH oxidation occurs at ca. 0.0 V vs SCE, reaching a well-defined peak at a disk peak peak potential, Edisk , ca. 0.2 V vs SCE. Plots of the disk current, idisk, at Edisk vs ω1/2 were linear with a close to zero intercept. Measurements in which Edisk of the rotating ring-disk electrode was scanned, while Ering was fixed at a value negative enough for solution phase NO to undergo reduction, yielded plots of iring vs Edisk, which mirrored the peak found for idisk and thus was consistent with NO being one of the predominant products of NH2OH oxidation. In situ infrared measurements provided evidence for N2O being produced at the same onset potential of NH2OH oxidation. The disk polarization curves could be reproduced by theoretical simulations involving an EEECE mechanism in which nitrite, one of the products of NH2OH oxidation, peak reacts with NH2OH yielding an electrochemically inert species. In accordance with theory, plots of idisk at Edisk as a function of [NH2OH] bent downward as [NH2OH] increased.

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INTRODUCTION Hydroxylamine, NH2OH, including its salts, is a chemical often used as a reducing agent in both organic and inorganic chemistry1 and also as an important intermediate in biological nitrification. Attention in the electrochemical area has been focused mostly on Pt electrodes in acidic and neutral electrolytes. According to Karabinas et al.,2 the oxidation of NH2OH in these media proceeds via formation of NOH adsorbed, NOH(ads), which undergoes dimerization to yield nitrous oxide, N2O, in solution and also reacts electrochemically at higher potentials generating NO. This conclusion was reached based on the analysis of online mass spectrometry data, which also showed a peak due to solution-phase NO2 in acidic, but not in neutral, media. It was also proposed by these authors that NOH(ads) could be oxidized directly on a Pt oxide surface to yield nitrite, a species that reacts homogeneously with NH2OH producing N2O, and oxidizes further at higher potentials to generate nitrate. More recently, Rosca et al. examined in more detail the same redox process under acidic conditions by both online MS and in situ IRAS3 and proposed based on theoretical calculations that the main intermediate is not NOH(ads), but NO(ads) produced via HNO(ads),4 which then acts as a poison for the further oxidation of NH2OH. Extension of their studies to include low index single crystal surfaces5 led these authors to conclude that the oxidation of NH2OH to NO appears to be a structure-insensitive process, although influenced by coadsorption of anions. © 2012 American Chemical Society

Our interests have focused on the use of Au electrodes for this process, a metal that displays no affinity for NO leading in principle to a much simpler reaction mechanism. Indeed voltammetric experiments performed in acetate buffer solution (pH 4) yielded two clearly defined peaks in the double layer region of Au.6 Furthermore, experiments involving rotating ring-disk electrode (RRDE) techniques afforded evidence that the two consecutive redox waves are associated primarily with oxidation of NH2OH to nitrite and nitrate, respectively. Analysis of the data collected with a rotating disk electrode (RDE) yielded nonlinear Levich plots pointing to complications in the overall mechanism. Insight into the possible involvement of adsorbed intermediates was obtained from in situ surface-enhanced Raman scattering (SERS), which revealed the presence of two distinct spectral features within the range in which NH2OH undergoes oxidation, centered at 803 cm−1 for 0.55 < E < 0.8 V and at 826 cm−1 for 1.0 < E < 1.40 V versus SCE, which were attributed, respectively, to NO2−(ads) and NO2(ads).7 This contribution presents Au(disk)|Au(ring) RRDE and in situ infrared reflection absorption spectroscopy (IRAS) data for the oxidation of NH2OH in phosphate buffer solutions at pH = 7. The rationale behind the selection of this specific electrolyte is 2-fold: (i) The chemicals involved are available Received: October 31, 2011 Revised: February 24, 2012 Published: February 24, 2012 6932

dx.doi.org/10.1021/jp2104566 | J. Phys. Chem. C 2012, 116, 6932−6942

The Journal of Physical Chemistry C

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RESULTS AND DISCUSSION Experimental. Shown in Figure 1 are cyclic voltammetric curves recorded under stagnant conditions in a 2 mM NH2OH

commercially in much higher purity than those employed in the earlier studies at lower pH; hence, any effects due to contaminants would be greatly reduced. (ii) There is a great body of literature addressing various aspects of the electrochemistry of simple nitrogen species of relevance to biological systems at physiological pH values.

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Article

EXPERIMENTAL SECTION

Electrochemistry. All measurements were performed with a rotating Au ring-Au disk electrode, Au|Au RRDE, (Pine Instruments, disk diameter, 0.457 cm; ring outer diameter, 0.538 cm; ring inner diameter, 0.493 cm; disk area, Adisk = 0.164 cm2; ring area, Aring = 0.037 cm2; gap, G = 180 μm; collection efficiency, N = 0.22) using a commercial rotator (Pine Instruments, model AFMSRX) and a bipotentiostat (Pine Instruments, model AFCBP1). Phosphate buffer (PB, pH 7) solutions were prepared from monosodium dihydrogen phosphate (Ultrapure, J.T. Baker) and disodium hydrogen phosphate (Ultrapure, J. T. Baker) and ultrapure water (18.3 MΩ cm, EASYpureUV system, Barnstead). Hydroxylamine (NH2OH, Aldrich, 99.999% pure) and NaNO2 (Fisher, ACS pure) were used without further purification. Nitric oxide (NO) was generated by the reaction between H2SO4 and NaNO2 and then bubbled through a wash bottle containing aqueous 3 M KOH to remove undesired gaseous products. Experiments were carried out in thoroughly deoxygenated solutions using a conventional all-glass, three-compartment cell with a carbon rod and a saturated calomel (SCE) as counter and reference electrodes, respectively. Spectroscopy. All in situ infrared spectra were acquired in the reflection−absorption configuration using a Bruker FTIR (Tensor 27) equipped with a narrow-band MCT detector (Infrared Associates) with p-polarized light at a 4 cm−1 resolution. The spectroelectrochemical cell employed in these experiments, described in detail elsewhere,8 incorporates a beveled CaF2 (60°) prism, an ellipsoidal polycrystalline Au foil encased in Kel-F (area = 1.7 cm2), and a thin Au foil as working and counter electrodes, respectively (see inset, Figure 6). The reference electrode was placed in a separate compartment connected to the main cell by means of a Teflon tubing to mitigate problems associated with contamination. During the actual spectral collection, the working electrode was pressed against the flat surface of the CaF2 window using a micrometer to reduce the thickness of the electrolyte film trapped within. Spectra representing the average of 500 consecutive interferometric scans were collected at fixed potentials, E, every 100 mV toward positive values starting at Eini = −0.4 V vs SCE, i.e., negative to the onset of NH2OH oxidation. Once the full acquisition was completed, the potential was scanned at 10 mV/s to the next value and the entire sequence repeated. The data are displayed as normalized potential difference spectra, i.e., −ΔR/R = [R(Eref) − R(E)]/R(Eref) vs wavenumber, cm−1, where R(Eref) is the IRAS spectrum obtained at Eini. Theoretical Simulations. All calculations were performed with COMSOL (Version 4.1) using the Transport of Diluted Species interface. The magnitudes of all the parameters, including the kinetic rate constants, were optimized to achieve best fits to the experimental data at one rotation rate, i.e., ω = 900 rpm, which were then used to predict results for other ω values.

Figure 1. Voltammetric curves recorded in a quiescent 2 mM NH2OH solution in 0.1 M PB (pH 7) at ν = 10 mV/s (red dotted curve) and 100 mV/s (green curve) over the potential range −0.9 ≤ E ≤ 0.65 V vs SCE. Also shown in this figure are voltammograms collected at ν = 10 mV/s over a much wider range, i.e., −0.9 ≤ E ≤ 1.25 V vs SCE in the same NH2OH solution (blue curve) and in neat PB electrolyte (black curve).

solution in 0.1 M phosphate buffer (PB, pH 7) at a scan rate, ν = 10 mV/s, over the potential ranges −0.9 ≤ E ≤ 0.65 V (red curve) and −0.9 ≤ E ≤ 1.3 V vs SCE (blue curve). Also displayed in this figure is the corresponding voltammogram obtained in the neat PB over the wider potential range (black curve). In analogy with the behavior found in acetate buffer (pH 4) solutions reported in our previous work, the oxidation of NH2OH on Au is characterized by two prominent waves, denoted as I and II, centered in this case at 0.05−0.1 V and ca. 0.75 V vs SCE. Unlike the data collected at pH 4, for which both peaks reached maximum values within the double layer region, the onset for peak II at pH 7 virtually coincides with the onset of Au oxide formation (see black curve in Figure 1). The small peak centered at ca. 0.5 V, labeled as I′ in the figure, may be attributed to the presence of solution phase NO, as a similar feature at about the same potential was observed in PB solutions containing dissolved NO (see red curve, Figure 2). Additional evidence in support of this view was obtained by increasing the scan rate, ν, an approach that enhances detection of soluble electroactive products, which over time diffuse away from the surface into the bulk solution. As shown in the green curve in Figure 1, two rather broad and shallow features were observed in NH2OH solutions during the scan toward negative potentials for voltammetric curves recorded at ν = 100 mV/s over the range −0.8 ≤ E ≤ −0.1 V vs SCE, in which NO undergoes reduction (see red curve, Figure 2).9 An indication that peak I′ is associated with the oxidation of NO was obtained from voltammetric experiments in which the scan toward positive potentials was reversed at a potential E+ in the range −0.1 to 0.6 V. Cursory inspection of the data collected (see Figure 3) revealed that the shallow features observed in the scan toward negative potentials become more intense as E+ was monotonically increased up to the onset of peak I′ and 6933



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Figure 4. Top Panel. Dynamic polarization curves (ν = 10 mV/s) for the Au disk of a Au|Au RRDE recorded in 0.1 M PB (pH 7) containing 2 mM NH2OH at ω = 100 (curve a), 400 (b), 900 (c), 1600 (d), and 2500 (e) rpm. Bottom Panel. Ring currents, iring, recorded with Ering = −0.7 V while scanning Edisk. The solid and dotted curves in these panels were obtained while scanning Edisk toward positive and negative values, respectively. Inset: Plot of ipeak for the scans toward positive potentials as a function of ω1/2. Linear Fit: slope, 0.0613 ± 8.1 × 10−4; intercept, 0.048 ± 0.027.

Figure 2. Voltammetric curves (ν = 10 mV/s) obtained in a quiescent 0.1 M PB (pH 7) solution containing either NO (red curve) over the range −0.9 ≤ E ≤ 0.65 vs SCE or 2 mM NaNO2 over the range −0.9 ≤ E ≤ 1.25 V vs SCE. The black curve is the corresponding voltammogram recorded in the neat electrolyte under otherwise identical conditions.

ν = 10 mV/s in 0.1 M PB (pH 7) containing 2 mM NH2OH for rotation rates ω = 100 (curve a), 400 (b), 900 (c), 1600 (d), and 2500 (e) rpm, where the solid and dotted lines correspond to scans toward positive and (subsequent) negative values, respectively. As indicated, the onset of NH2OH oxidation for all these curves occurs at ca. −50 mV vs SCE followed by a rapid peak , in the range increase in the current reaching a maximum, Edisk 0.15−0.25 V vs SCE. Thereafter, the current decreases down to a minimum at ca. 0.5 V and later increases up to the onset of Au oxide formation in solutions devoid of NH2OH. A plot of the peak , ipeak, as a function of ω1/2 yielded a straight line current at Edisk with a close to zero intercept (intercept, 0.048 ± 0.02701; and slope, 0.0613 ± 8.144 × 10−4) as would be expected for a diffusion-limited process (see inset, Figure 4). The ring currents, iring, recorded while scanning Edisk are displayed in the bottom panel in this figure. As clearly evidenced by these data, the behavior of iring mirrors that found for idisk up to Edisk ca. 0.5 V, which suggests that the amount of NO being produced at the disk is proportional to idisk; however, for Edisk > 0.5 V, iring decreases. It should be stressed in this regard that nitrite, nitrate, nitrous oxide, and dinitrogen, which are some of the possible products of NH2OH oxidation, are electrochemically inactive on Au in this potential region and thus cannot account for the observed currents. In particular, no NH2OH could be detected in Au|Au RRDE experiments performed in nitrite containing PB solutions in which the Au ring was polarized at a potential of 0.1 V vs SCE while scanning Edisk toward negative values down to the onset of hydrogen evolution (see Supporting Information). Dynamic polarization curves were also recorded for the oxidation of NH2OH as a function of [NH2OH] in the range 0.25 ≤ [NH2OH] ≤ 2.5 mM (see upper panel, Figure 5) with the ring polarized at Ering = −0.7 V (see lower panel in the same figure) and at ω = 900 rpm, yielding, as expected, results qualitatively very similar to those shown in Figure 4. In analogy with the behavior found in acetate buffer (pH = 4) in our earlier publication, a plot of ipeak vs [NH2OH] was found to

Figure 3. Series of voltammetric curves recorded at ν = 100 mV/s in the same solution as that specified in the caption of Figure 1 for values of E+ (see text) = −0.1 (red trace); 0.03 (green); 0.2 (blue); 0.4 (light blue); and 0.6 (magenta) V vs Ag|AgCl. Inset: Expanded curves along the current axis in the double layer region.

then decreased (see magenta curve for E+ = 0.6 V and expanded view in the inset) as E+ was further increased in accordance with the suggestion above. It is important to stress that nitrite does not undergo reduction on Au within this potential range (see green curve in Figure 2 and Figure S1 in the Supporting Information) but oxidizes readily at E > 0.65 V, i.e., very close to the fast rising current observed in NO containing solutions (see red curve, Figure 2). Some aspects of this latter process will be addressed below. On the basis of these observations, a series of Au|Au RRDE experiments were performed aimed at correlating the amount of NO produced during the oxidation of NH2OH as a function of the applied disk potential, Edisk. For these measurements, the potential of the ring, Ering, was set at −0.7 V, i.e., negative enough for NO reduction to ensue. Shown in the upper panel of Figure 4 are dynamic polarization curves recorded with the Au disk at 6934



Article peak i.e., Edisk , are significantly lower than those expected for the reduction of the product generated at the disk assuming the two reactions are precisely the opposite of one another, e.g., NH2OH oxidation to NO on the disk and NO reduction to NH2OH on the ring. This may not be surprising, as preliminary data collected in our laboratory have afforded evidence that the reduction of solution phase NO in aqueous electrolytes does not proceed under pure diffusion control over the entire potential range accessible in this electrolyte. Yet another factor that could contribute to this disparity is the possibility that NH2OH oxidation generates, in addition to NO, one or more products which are not electrochemically active down to the onset of hydrogen evolution on the Au ring. Insight into the nature of these possible products was gained from in situ infrared reflection absorption spectroscopy (IRAS) measurements collected in the same solution as that specified in the figure captions above at fixed potentials over the range −0.4 < E < 1.1 V vs SCE in increments of 0.1 V. Shown in the left panel of Figure 6 are a series of in situ normalized IRAS spectra over the region 2275−2100 cm−1 as a function of the applied potential, using the spectrum at E = −0.4 V as a reference, where all the traces have been displaced along the y-axis for clarity. Also eliminated were contributions due to the featureless sloping background by simply subtracting from each of the resulting spectra collected at E = −0.2 V, yielding a prominent peak centered at 2230 cm−1, characteristic of nitrous oxide, N2O. These data were used to construct a plot of the integrated intensity of the peak, I2230, vs E (see right panel in this figure). As clearly evidenced by the results, the onset of N2O formation, i.e., ca. 0.5 V vs SCE, virtually coincides with the onset potential for NH2OH oxidation. In addition, I2230 reaches a maximum at ca. 0.3 V, which is very

Figure 5. Top Panel. Dynamic polarization curves (ν = 10 mV/s) for the Au disk of a Au|Au RRDE recorded in 0.1 M PB (pH 7) as a function of [NH2OH] in the range 0.25 (bottom curve) to 2.5 mM (top curve) in 0.25 mM increments at ω = 900 rpm. Bottom Panel. Ring currents, iring, recorded with the ring polarized at Ering = −0.7 V, while scanning Edisk. Inset: Plot of ipeak vs [NH2OH] (scattered symbols). The smooth straight line was added to show that the experimental data deviate from linearity.

bend downward as [NH2OH] increased (see inset). A model that accounts quantitatively for this peculiar phenomenon, including the shape of the disk polarization curve, will be presented later in this work. Careful analysis of the RRDE data in Figures 4 and 5 indicates that the actual values of iring recorded at the peak maximum,

Figure 6. Left Panel. Series of in situ FTIR spectra at fixed potentials over the range −0.1 < E < 1.1 V vs SCE in steps of 0.1 V in the sequence specified by the arrow in the same solution as that used in previous experiments. Right Panel. Plot of the integrated intensity of the prominent feature at 2230 cm−1 as a function of E based on the data in the Left Panel. Inset: Schematic diagram of the cell for in situ IRAS measurements. (A) Micrometer, (B) counter electrode, (C) tubing to reference electrode, (D) Teflon lid, (E) glass body, (F) O-ring, (G) gold working electrode, (H) CaF2 prism. 6935



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close to that found for the peak in the dynamic polarization curves above. Since N2O cannot be oxidized on Au, the decrease in the intensity of the spectral feature as the potential is increased can be ascribed to the loss of this species due to diffusion away from the thin electrolyte layer between the electrode and the window and thus beyond the path of the IR beam. Yet another aspect that requires attention is the effect of impurities on the polarization curves. Although no systematic study was undertaken, differences were found between data collected as a function of the negative potential limit, E− lim. Specifically, the onset potential for NH2OH oxidation for E−lim = −0.7 V (see red and black curves in Figure 7, where solid and

Figure 8. Dynamic polarization curves recorded with a polycrystalline Au RDE (black) and a polycrystalline Pt RDE (red) at ω = 900 rpm in 0.1 M PB (pH 7) containing 2 mM NH2OH.

generation of solution phase nitrous acid, NO2−, which then reacts homogeneously with NH2OH to yield nitrous oxide, N2O. As shown in Appendix A, however, this EC-type mechanism fails to reproduce the peak observed for the oxidation of NH2OH on Au in 0.1 M phosphate buffer pH 7 making it necessary to introduce further complexities to the reaction sequence. The mechanism herein proposed, specified in Scheme 1, is based on electrochemical and spectroscopic data presented

Figure 7. Dynamic polarization curves (ν = 10 mV/s) recorded with the Au disk of a Au|Au RRDE recorded in 0.1 M PB (pH 7) containing 2 mM NH2OH at ω = 900 rpm over region I, i.e., −0.7 ≤ Edisk ≤ 0.65 V vs SCE (black curve). The green curve represents the sixth scan after cycling the electrode over region II, i.e., −0.3 ≤ Edisk ≤ 0.65 V vs SCE. The red curve was obtained after the negative potential limit was extended down to −0.7 V vs SCE. Solid and dotted curves were recorded for scans toward positive and negative potentials, respectively.

Scheme 1

dotted lines represent scans toward positive and negative potentials) was more negative than that found for E−lim = −0.3 V (see green curves in the same figure). This behavior is consistent with the desorption of species capable of adsorbing at high potentials, such as anions (in all likelihood chloride in this case) which would desorb at sufficiently negative potentials and thus open surface sites for the process to ensue. Finally, it is of interest to compare the behavior observed for polycrystalline Au with that for polycrystalline Pt. As shown in Figure 8, the curves found upon reversing the scan at 1.0 V were very similar for these two metals. However, as reported by Karabinas et al.,2 the currents for Pt in the potential range 0−0.5 V during the forward scan were much smaller than those observed during the subsequent scan in the negative direction. This hysteresis is attributed to the adsorption of NO on Pt a process that does not occur on Au and suggests that the oxidation of NH2OH may not be a surface-specific process provided the Pt and Au surfaces are devoid of strongly adsorbed species.10 Studies involving single crystals are currently underway to resolve some of these issues. Theoretical. Our initial emphasis was focused on an electrochemical-chemical (EC) mechanism suggested originally by Piela et al.11 for the oxidation of NH2OH on Pt in acid media. According to these authors, the first step in the process involves

above and additional information reported in the literature.12 As indicated in the scheme, it involves four electrochemical processes denoted as a, b, c, and e and a single homogeneous chemical reaction, d, where all species are assumed to be in the solution phase. In agreement with experimental observations, NH2OH can be oxidized to yield via parallel pathways nitric oxide, NO, and N2O (steps a and e in Scheme 1, respectively). Unlike N2O, which is electrochemically inert on Au,10 NO can undergo oxidation generating NO2− (b), which, as mentioned above, reacts with NH2OH in the solution phase (d) to produce N2O and also further oxidizes to nitrate (c). The potential-dependent rate constants for the heterogeneous electron transfer reactions, assumed for simplicity to be first order in the reactant, are denoted as ki (i = a, b, c, and e in Scheme 1), whereas the rate constant for the homogeneous reaction, d, for which the order will be specified later in this work is defined as k‴. All redox reactions will be assumed to be (kinetically) irreversible and to follow Butler−Volmer kinetics, with rate constants given by ki = k0io = exp[(αiniFE)/(RT)], 6936



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i = a, b, c, e, where k0io = k0io exp[(αiniFEi0)/(RT)] represents the first-order potential-independent standard heterogeneous rate constant, αi the transfer coefficient, and Ei0 the (Nernstian) reversible potential of reaction i, respectively. From a mathematical viewpoint, the problem so stated reduces to finding solutions for the three coupled differential equations that govern convective diffusion to a RDE for solution phase NH2OH, NO, and NO2− (see eqs I.1, I.2, and I.3 in Table 1), subject to the boundary conditions specified in Table 2. The symbols cJ and cJo in these tables represent the concentrations of species J = NH2OH, NO, and NO2− in the solution and at the surface of the electrode, respectively, and DJ represents their corresponding diffusion coefficient. Furthermore, uy = (ων)1/2H(γ) is the fluid velocity along y, the axis normal to the electrode surface, where ω is the rotation rate of the electrode, ν the kinematic viscosity of the solution, and H(γ) = −[aγ2 − (1/3)γ3 − (1/6)bγ4 − (1/30)b2γ5], where γ = (ω/ν)1/2y is a dimensionless axial distance.13 Also shown in Table 1 are the corresponding differential equations written in terms of dimensionless concentrations, i.e., C NH 2OH = ∞ ∞ )]; CNO = [(cNO)/(cNH )]; and CNO2− = [(cNH2OH)/(cNH 2OH 2OH ∞ ∞ [(cNO2−)/(cNH2OH )], where cNH2OH is the bulk concentration of NH2OH and of dimensionless parameters, i.e., the Schmidt ∞ numbers of species J, SJ, and m = [(cNH )2k‴]/ω. The term 2OH 2 − k‴cNH2OHcNO2 in eqs I.1 and I.3 accounts for the rate of the homogeneous reaction between NH2OH and NO2− (see d in Scheme 1) and was selected based on kinetic studies reported by Döring and Gehlen12 (see Appendix B). Finally, the parameter κiJ = [(kiSJ1/2)/((DJω)1/2)] (i = a,b,c,e; J = NH2OH, NO, NO2−) in Table 2 is a dimensionless, potential-dependent heterogeneous rate constant. Within this framework, the total current density can be expressed in terms of contributions due to the four electrochemical reactions in Scheme 1, namely

0 0 0 − + 2k c 0 i = (3kac NH + k bc NO + 2kcc NO e NH2OH)F 2OH 2 (6)

On this basis, the dimensionless total current I can be written in terms of dimensionless quantities as follows −

0 0 NO2 0 I = (3κ aHAMCNH + κ NO CNO − b CNO + 2κ c 2OH 2 0 + 2κ eNH2OHCNH ) 2OH

(7)

For all the simulations, the temperature T was always fixed at 298 K. Although no differences in the results were observed for γ > 0.3, all calculations were performed setting γ = 1.0, which, assuming ν = 0.01 cm2 s−1 and the value ω above, corresponds to y ca. 0.01 cm. Shown in Figure 9 are plots of I as a function of the dimensionless applied potential, FE/RT, for values of m = 0, i.e. no homogeneous reaction (green), m = 0.025 (black), 0.25 (red), 2.5 (blue), and 25.0 (magenta), for simulations involving selected values of the parameters involved, as specified in the caption in this figure. As indicated, all curves virtually overlap both for small potentials, EF/RT < 7, a region where the predominant reaction is the oxidation of NH2OH to NO (a in Scheme 1) and N2O (e), and for large potentials, EF/RT > ca. 25, where nitrate generation dominates (c). For intermediate values, however, the shape of the curves depends markedly on the value of m. Specifically, the smooth and monotonic increase in the current between the two plateaus observed for very small m gives rise to a dip that becomes more pronounced as m increases. Since m is inversely proportional to ω, the behavior observed mimics that found experimentally (vide infra). A more detailed understanding of this peculiar effect can be gleaned from the analysis of the changes in the contributions of the partial currents associated with each of the faradaic processes to the total dimensionless current, I, induced by the

Table 1. Governing Differential Equations dimensioned variables 2

D NH2OH D NO

∂ c NH2OH ∂y 2

∂ 2c NO ∂y 2

D NO2−

− uy

∂c NH2OH

2 −=0 − k‴c NH2OHc NO

∂y

2

∂c − uy NO = 0 ∂y

∂ 2c NO2− ∂y 2

− uy

dimensionless variables

eq

∂ 2CNH2OH ∂CNH2OH 2 −=0 − S NH2OHH(γ) − S NH2OHmCNH2OHCNO 2 2 ∂γ ∂γ

I.1

∂ 2CNO

I.2

∂γ2

∂c NO2− ∂y

2 −=0 − k‴c NH2OHc NO 2

− SNOH(γ)

∂ 2CNO2− ∂γ2

∂CNO =0 ∂γ

− S NO2−H(γ)

∂CNO2− ∂γ

I.3 2 −=0 − S NO2−mCNH2OHCNO 2

Table 2. Boundary Conditions boundary y = ∞, γ = ∞

dimensioned variables


∞ c NH2OH = c NH ;c = 0; c NO2− = 0 2OH NO

y = 0, γ = 0

− D NH2OH

− D NO2−

∂c NH2OH ∂y

∂c − D NO NO ∂y

y=0

0 0 = − kac NH − kec NH 2OH 2OH

−

∂CNH2OH

−

∂CNO2−

y=0

∂c NO2− ∂y

CNH2OH = 1; CNO = 0;

y=0

0 0 − = k bc NO − kcc NO 2

0 0 = kac NH − k bc NO 2OH

−

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∂γ

∂γ

CNO2−

=0

0 0 = −κ aNH2OHCNH − κ eNH2OHCNH 2OH 2OH

eq II.1 II.2

γ= 0

γ= 0

0 NO2−C 0 = κ NO b CNO − κ c NO2−

∂CNO 0 0 = κ aNH2OHCNH − κ NO b CNO 2OH ∂γ γ= 0

II.3

II.4



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dimensionless fluxes of the electroactive species involved, −[(∂CJ)/(∂γ)]|γ=0, for J = NH2OH, NO, and NO2− in Panels A and B of Figure 11, respectively. As indicated in Panel A,

∞ Figure 9. Plots of I vs FE/RT for various values of m = (cNH )2 2OH k‴/ω = 0 (green), 0.025 (black), 0.25 (red), 2.5 (blue), and o o = 3.19 × 10−5 cm s−1, k0b = 2.07 × 25.0 (magenta), assuming k0a o o = 4.90 × 10−12 cm s−1, k0e = 2.22 × 10−4 cm s−1, and 10−6 cm s−1, k0c SNH2OH = SNO = SNO2− = 3.33 × 103.

Figure 11. Dimensionless surface concentrations (Panel A) and dimensionless surface fluxes (Panel B) for NH2OH, NO, and NO2− as a function of EF/RT using the same set of parameters specified in the caption of Figure 1 for m = 0 (solid circles) and m = 25 (open circles). 0 drops to a very small value at a potential negative to the CNH 2OH onset of dip and remains virtually zero for larger potentials as for a conventional diffusion limited process. Not surprisingly, the corresponding surface concentrations of NO and NO2− display peak-like shapes which vanish at sufficiently positive potentials as diffusion limited conditions are reached. The surface fluxes of these two species were also found to display peaklike shapes (see Panel B, Figure 11) and also decreased to negligible values at high enough potentials. The surface flux associated with NH2OH, i.e., −[(∂CNH2OH)/(∂γ)]|γ=0 (see eq II.2 in Table 2), first increases with potential reaching a maximum at EF/RT of ca. 8 and then decreases in the region of dip (see Figure 10) and, subsequently, increases for higher potentials (see Panel B, Figure 11) to attain a limiting value. In fact, the decrease in −[(∂CNH2OH)/(∂γ)]|γ=0 induced by reaction d becomes evident in Figure 12, which compares the concentration profiles of NH2OH within the diffusion boundary layer for m = 0 (red) and m = 25 (black). As indicated therein, the homogeneous reaction elicits a decrease in the overall values of CNH2OH, leading to a corresponding decrease in the magnitude of the flux at the interface. A quantitative account of the experimental data based on this proposed model could be obtained by optimizing the values of the parameters involved using the data collected at ω = 900 rpm (see Table 3) as a basis. Shown in Figure 13 is a comparison between the measured (solid lines) and the best fit polarization curves for five different rotation rates calculated using the parameters listed in Table 3 (scattered symbols). The much lower currents observed experimentally for the highest rotation rate are in all likelihood due to the gradual accumulation of adsorbed anionic impurities on the surface, in all likelihood chloride, at potentials positive to the point of zero charge, for which the rate would increase as ω increases. This effect would also be responsible for the lack of linearity of the Levich plots, i.e., ipeak vs ω1/2, in Figure 14 observed in our

Figure 10. Plots of the total, I, and partial dimensionless currents for each of the redox processes in Scheme 1 (see labels) vs FE/RT for ∞ values of m = (cNH )2k‴/ω = 0 (solid symbols) and m = 25 (open 2OH symbols). The values of all other parameters are specified in the caption for Figure 1. The partial current due to reaction e is not affected by the homogeneous process.

homogeneous process shown in Figure 10. As clearly evident from a comparison between the curves for m = 0 (solid symbols) and m = 25 (open symbols) therein, the dip in I in the range 10 < EF/RT < 23 is derived from a corresponding dip in the partial current due to reaction a (magenta) induced by reaction d. In fact, the latter is also responsible for a shift in the onset of reaction b toward higher dimensionless potentials. Finally, as could have been expected, the contribution due to reaction e to I does not seem to be affected by the homogeneous reaction. Further insight into the origin of the dip can be gained from plots of the dimensionless surface concentrations CJ0 and 6938



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Figure 12. Dimensionless concentration profiles of NH2OH using the set of parameters specified in the caption of Figure 1 for m = 0 (red) and m = 25 (black) for EF/RT = 20. Figure 14. Comparison between the experimental (open circles) and theoretical (solid circles) peak currents as a function of ω1/2 based on the data in Figure 5. A linear fit to the simulated data yielded a slope S = 0.0661 ± 4.62 × 10 −4, an intercept I = −0.0818 ± 0.01532, and R2 = 0.9998.

Table 3. Best Fit Parameters for the Mechanism in Scheme 1 koa kob koc koe

8 5 2 5

× × × ×

DNH2OH

10−3 cm s−1 10−3 cm s−1 10−4 cm s−1 10−3 cm s−1 = DNO = DNO2− k‴ ω

Ea0 Eb0 Ec0 Ee0

0.10 0.31 0.54 0.10

V αa 0.3 V αb 0.5 V αc 0.24 V αe 0.3 4.7 × 10−6 cm2 s−1

6.8 × 1014 cm6 mol−2 s−1 900 rpm

Figure 15. Comparison of the peak potential, Epeak, vs ω based on the experimental data in Figure 5 (open circles) and on the best fit parameters for the proposed model (solid circles).

Figure 13. Plots of experimental (solid curves, see Figure 4) and simulated polarization curves (scattered symbols) using the best fit parameters listed in Table 3 for the same rotation rates as specified in Figure 4.

Although the mechanism herein presented indeed accounts for the experimental results, its uniqueness cannot be ascertained. Some aspects of this issue were examined by introducing modifications to the model. The first involved a change in the rate law that governs the homogeneous reaction (d) in Scheme 1. Specifically, calculations were performed assuming rather arbitrarily that the process is first order in each of the two reactants. As shown in Appendix C, reasonable parameters could be found that reproduced the experimental data as accurately as the more complex rate law employed originally.

laboratory for acetate buffer solutions pH 46 and also for the slightly smaller Levich slope than that which would be predicted theoretically for a three-electron transfer reaction. Within these limitations, which would amount to a difference of ca. 5%, a Au ring polarized at the peak potential could therefore be used to detect quantitatively hydroxylamine at the disk in a conventional RRDE experiment carried out in this media. Further support for this proposed model is provided by the predicted shift in the peak potential, Epeak, toward higher values as ω increases (see solid circles in the same figure), which is in very good agreement with the behavior found experimentally (see open circles, Figure 15).

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CONCLUDING REMARKS The electrochemical and spectroscopic data presented in this paper for the oxidation of NH2OH in phosphate buffer pH 7 have been found to be consistent with a mechanism that 6939



Article

Table A.1. Governing Differential Equations dimensioned Variables 2

D NH2OH

D NO2−

∂ c NH2OH ∂y 2

∂ 2c NO2− ∂y 2

− uy

− uy

∂c NH2OH ∂y

∂c NO2− ∂y

2 −=0 − k‴c NH2OHc NO 2

2 −=0 − k‴c NH2OHc NO


eq

∂ 2CNH2OH ∂CNH2OH 2 −=0 − S NH2OHH(γ) − S NH2OHmCNH2OHCNO 2 ∂γ ∂γ2

A.I

∂ 2CNO2−

A.II

∂γ2

2

− S NO2−H(γ)

∂CNO2−

2 −=0 − S NO2−mCNH2OHCNO

∂γ

2

Table A.2. Boundary Conditions boundary

dimensioned variables

y = ∞, γ = ∞


∞ −= 0 c NH2OH = c NH ;c 2OH NO2

y = 0, γ = 0

− D NH2OH

− D NO2

∂c NH2OH ∂y

CNH2OH = 1; CNO2 = 0

0 = − k f c NH 2OH

0 = k f c NH 2OH

−

∂γ ∂CNO2 ∂γ

NH2OH 0 CNH2OH

eq A.III A.IV

= −κ f γ= 0

−

NH2OH 0 CNH2OH

A.V

= κf γ= 0

k‴

NO2− + NH2OH ⎯→ ⎯ N2O + H2O + OH−

(A.2)

The corresponding governing differential equations and boundary conditions are given in Tables A.1 and A.2, respectively, where the rate law for the homogeneous reaction (A.2) was assumed to be third order as reported in the literature (see Appendix B). The total current in this case is due to the electrochemical reaction A.1, namely,

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APPENDIX A The primary objective of this Appendix is to examine whether a simple EC mechanism of the type shown in eqs A.1and A.2 below can reproduce the unique shape of the polarization curves observed for NH2OH oxidation on a Au RDE in phosphate buffer pH 7. As indicated therein, the reaction sequence involves formation of nitrous acid, NO2−, via an initial electrochemical process (A.1), which then reacts with NH2OH in solution to yield nitrous oxide (A.2). NH2OH + H2O → NO2− + 4e + 5H+

∂CNH2OH

y=0

involves generation of solution phase NO and HNO 2 (or nitrite) at small overpotentials, i.e., whereby the latter reacts homogeneously with NH2OH to yield N2O and, at higher potentials, undergoes full oxidation generating nitrate. Support for this mechanism has been obtained from theoretical simulations.

kf

−

y=0

∂c NO2− − ∂y

−

0 i = n f Fk f c NH 2OH

(A.3)

which can be expressed in terms of the dimensionless quantities selected as follows NH2OH 0 CNH2OH

I = 4κ f

(A.1)

(A.4)

Shown in Figure A.1 are simulated dimensionless polarization curves using parameter values for the constants involved given in the caption. As clearly indicated, the following chemical reaction in this simple EC mechanism does not elicit a peak and as such does not adequately describe the experimentally observed results.

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APPENDIX B According to Döring and Gehlen,12 the rate law that governs the homogeneous reaction d in Scheme 1 in acetate solutions is given by ν=

k[HNO2 ]2 [NH2OH][Ac−] k1o[Ac−] + (αo + α1[NO2−])[NH2OH]

(B.1)

Table B.1 parameters Reported by Döring and Gehlen12 α1 = 219.4 s αo = 2.167 × 10−3 s mol cm−3 (k1″)/(k″) = 3.4 × 10−6 mol2 s cm−6

∞ FE/RT for various values of m = (cNH )2k‴/ω, 2OH 2.5 (magenta), and 25.0 (green), assuming k0fo =

Figure A.1. Plots of I vs i.e., 0 (red), 0.25 (blue), 7.47 × 10−5 in cm/s and SNH2OH = SNO2− = 2128.

6940

parameters evaluated using the raw data reported by Döring and Gehlen12 α1 = 846.23 s αo = 2.685 × 10−3 s mol cm−3 (k1″)/(k″) = 0.75 × 10−6 mol2 s cm−6



Article

Table C.1. Governing Differential Equations dimensioned variables


eq

Model I

D NH2OH D NO

∂ 2c NH2OH ∂y 2

− uy

∂c NH2OH ∂y

− k″c NH2OHc NO2− = 0

NO − u ∂c NO = 0 y ∂y ∂y 2

∂ 2c NO2− ∂y

2

− uy

∂γ2

− S NH2OHH(γ)

∂CNH2OH ∂γ

C.I

− SNH2OHmC NH2OHCNO2− = 0 C.II

∂ 2CNO ∂C − S NOH(γ) NO = 0 ∂γ ∂γ2

∂ 2c

D NO2−

∂ 2CNH2OH

∂c NO2− ∂y

− k″c NH2OHc NO2− = 0

∂ 2CNO2− ∂γ2

− SNO2−H(γ)

C.III

∂CNO2−

− SNO2−mC NH2OHCNO2− = 0

∂γ

Model II

D NH2OH D NO

∂ 2c NH2OH ∂y 2

− uy

∂c NH2OH ∂y

2 −= 0 c − k IIIc NH 2OH NO2

NO − u ∂c NO = 0 y ∂y ∂y 2

∂ 2c NO2− ∂y 2

− uy

∂c NO2− ∂y

∂γ2

− S NH2OHH(γ)

2

∂ 2c

D NO2−

∂ 2CNH2OH

∂ CNO ∂γ2 2 −= 0 − k IIIc NH c 2OH NO2

− S NOH(γ)

∂ 2CNO2− ∂γ2

∂CNH2OH ∂γ

2 −= 0 − SNH2OHmC NH C 2OH NO2

C.V

∂CNO =0 ∂γ

− S NO2−H(γ)

∂CNO2− ∂γ

C.IV

2 −= 0 C − S NO2−mC NH 2OH NO2

C.VI

Figure C.1. Comparison between the experimental (scattered black symbols) and best fit simulated polarization curves (solid lines) based on the second-order model (Left Panel, Di = 4.8 × 10−6 cm2/s, k″ = 1.5 × 108 cm3/mol s (red symbols), k″ = 0 cm3/mol s (blue)) and third-order model (Right Panel, Di = 4.9 × 10−6 cm2/s, kIII = 8.8 × 108 cm6/mol2 s (red), kIII = 0 cm6/mol2 s (blue)) described in this Appendix for ω = 900 rpm. All other best fitting parameters are listed in Table C.2.

which for k1o[Ac−] ≫ (αo + α1[NO2−])[NO2OH]+ reduces to ν = k‴[HNO2 ]2 [NH2OH]

Note that the expression given in the Abstract of this reference has a typographical error, which upon rearrangement can be expressed as

(B.2)

Our analysis assumes that the phosphate species play the same role as acetate in their studies and that the expression in eq B.2 would still be valid at the higher pH values employed in our studies. Although this assumption may be subject to question, the fact that the overall features do not seem to be very sensitive to the rate law (as illustrated in Appendix C) lends some credence to our approach.

[HNO2 ]o2 [NH2OH]o k″ k″ = 1 + 2 [HNO2 ]o [NH2OH]o νo k″ k″ (B.4)

where ν is the rate and the subscript o denotes initial values for all quantities involved. On this basis, a plot of {([HNO2]o2)([NH2OH]o)}/νo vs [NH2OH]0 for a fixed [HNO2]o should be linear; hence, by selecting various values of [HNO2]o, the resulting straight lines should yield a common intercept, i.e., (k1″/k″) in eq B.4. If the slopes of these lines, namely, (k2″/k″)[HNO2]o, are then plotted against the various [HNO2]o, the intercept and slope of the resulting straight line define the values of the parameters αo and α1 in eq B.1. Although these constants were not employed in the rate expression employed in our work, i.e., B.2,

On the Values of the Constants αo and α1 in Equation B.1: A Critical Assessment

The rate law proposed by Döring and Gehlen12 (see eq B.1 above) was based on the analysis of initial rate data collected in their laboratory, which yielded the following empirical expression νo =

k″[HNO2 ]o2 [NH2OH]o k1″ + k2″[HNO2 ]o [NH2OH]o

(B.3) 6941



Article

digitization of the authors’ data and subsequent evaluation of the respective constants resulted in values different from those reported by these authors, as shown in Table B.1 below.

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APPENDIX C

To be examined in this section is the same mechanism as that specified in Scheme 1 assuming the homogeneous reaction follows either a second-order rate law, first order in NH2OH and NO2− (Model I), or a third-order rate law, second order in NH2OH and first order in NO2− (Model II). The governing differential equations are given in Table C.1 below where k″ is ∞ )/ω and kIII is the second-order rate constant and m = (k″cNH 2OH III ∞ the third-order rate constant and m = (k cNH2OH )2/ω. The boundary conditions are the same as those specified in Table 2. The main objective of this exercise is to determine whether a different rate law could also account for the results observed experimentally with the RDE. The total current has contributions due to the four electrochemical reactions, as specified in dimensioned and dimensionless variables in eqs 6 and 7, respectively. As clearly evident from the data shown in Figure C.1, reasonable parameters can be found that fit the experimental data with a second- and also a third-order rate law of the type assumed herein.

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8 5 2 5

× × × ×

10−3 cm 10−3 cm 10−4 cm 10−3 cm

s−1 s−1 s−1 s−1

Ea0 Eb0 Ec0 Ee0

0.10 0.31 0.54 0.10

V V V V

αa αb αc αe

0.3 0.5 0.24 0.3

ASSOCIATED CONTENT

S Supporting Information *

Dynamic polarization curves for the Au disk electrode of a Au| Au RRDE recorded in 0.1 M PB (pH 7) containing 2 mM NaNO2 at 400 rpm and ring currents, iring, recorded with Ering = 0.2 V, while scanning Edisk. This material is available free of charge via the Internet at http://pubs.acs.org.

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REFERENCES

(1) Tauszik, G. R.; Crocetta, P. Appl. Catal. 1985, 17, 1−21. (2) Karabinas, P.; Wolter, O.; Heitbaum, J. Ber. Bunsen-Ges. Phys. Chem. Chem. Phys. 1984, 88, 1191−1196. (3) Rosca, V.; Beltramo, G. L.; Koper, M. T. M. J. Electroanal. Chem. 2004, 566 (1), 53−62. (4) de Vooys, A. C. A.; Koper, M. T. M.; van Santen, R. A.; van Veen, J. A. R. J. Catal. 2001, 202, 387−394. (5) Rosca, V.; Beltramo, G. L.; Koper, M. T. M. J. Phys. Chem. B 2004, 108, 8294−8304. (6) Chen, Y.; de Godoi, D. R. M.; Scherson, D. J. Electrochem. Soc. 2011, 158, F29−F35. (7) Godoi, D. R. M.; Chen, Y.; Zhu, H.; Scherson, D. Langmuir 2011, 26, 15711−15713. (8) Huang, H.; Zhao, M.; Xing, X.; Bae, I. T.; Scherson, D. J. Electroanal. Chem. 1990, 293, 279−284. (9) Suzuki, S.; Nakato, T.; Hattori, H.; Kita, H. J. Electroanal. Chem. 1995, 396, 143−150. (10) Rosca, V.; Duca, M.; de Groot, M. T.; Koper, M. T. M. Chem. Rev. 2009, 109, 2209−2244. (11) Piela, B.; Wrona, P. K. J. Electrochem. Soc. 2004, 151, E69−E79. (12) Döring, C.; Gehlen, H. Z. Anorg. Chem. 1961, 312, 32−44. (13) Bard, A. J.; Faulkner, L. R. Electrochemical Methods: Fundamentals and Applications; Wiley: New York, 1980.

Table C.2. Best Fit Parameters for Model I and Model II koa kob koc koe

k‴ = third-order reaction rate constant for the homogeneous reaction between NH2OH and NO2− in cm6/mol2 s. Second order in NO2− k″ = second-order reaction rate constant for the homogeneous reaction between NH2OH and NO2− in cm3/mol s. First order in each of the species kIII = third-order reaction rate constant for the homogeneous reaction between NH2OH and NO2− in cm6/mol2 s. Second order in NH2OH y = axial distance in cm uy = velocity in the axial direction in cm/s ω = angular velocity in rad s−1 ν = kinematic viscosity in cm2/s αi = transfer coefficient of reaction i γ = a dimensionless axial distance SJ = Schmidt number of species J m = a dimensionless parameter

AUTHOR INFORMATION

Corresponding Author

*Phone: 216 368 5186. Fax: 216 368 3006. E-mail: daniel. [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the National Science Foundation (CHE - 0911621). LIST OF SYMBOLS DJ = diffusion coefficient of species J in cm2/s, J = NH2OH, NO, NO2− cJ = concentration of species J in mol/cm3 cJ∞ = bulk concentration of the reactant J in mol/cm3 CJ = dimensionless concentration of species J ki = first-order heterogeneous rate constant for reaction i = a, b, c, e in cm s−1 ni = number of electrons transferred in each of the heterogeneous electron transfer reactions 6942


Oxidation of Hydroxylamine on Gold Electrodes in Aqueous

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