Article
pubs.acs.org/JPCC

Electrochemical Reaction of Aqueous Iron Sulfate Solutions Studied
by Fe L‑Edge Soft X‑ray Absorption Spectroscopy
Masanari Nagasaka,*,†,‡ Hayato Yuzawa,† Toshio Horigome,† Adam P. Hitchcock,§and Nobuhiro Kosugi†,‡†

Institute for Molecular Science, Myodaiji, Okazaki 444-8585, JapanThe Graduate University for Advanced Studies, Myodaiji, Okazaki 444-8585, Japan§
Department of Chemistry and Chemical Biology, McMaster University, Hamilton, ON L8S 4M1, Canada‡

ABSTRACT: Change in valence of Fe ions in aqueous ironsulfate solutions at diﬀerent potentials has been studied by Fe
L-edge soft X-ray absorption spectroscopy (XAS) in transmission mode. Each XAS spectrum is measured at a constantpotential by using a liquid cell with built-in electrodes. A
nonlinear oxidation of Fe(II) ions to Fe(III) ions is observed
when the potential is increased from 0.0 to 0.9 V. Twoprocesses are found in the oxidation: one is a simple oxidation
process and the other is a process involving the sulfate ions.
The potential peak in the latter process is changed with
diﬀerent scanning rates because the sulfate ions aﬀect electrode kinetic parameters and diﬀusion coeﬃcients. The reduction of
Fe(III) ions to Fe(II) ions shows a linear proﬁle when the potential is decreased from 0.9 to −0.4 V. The mechanism of these Fe
redox processes is discussed by correlating the XAS results with cyclic voltammetry results at diﬀerent scanning rates.

1. INTRODUCTION
To get a better understanding of electrochemistry, it isnecessary to know structures of electrolytes including electric
double layers at diﬀerent potentials. The structures of electric
double layers were mainly studied by using electrolyte
molecules adsorbed on single-crystal electrode surfaces undervacuum. For example, the structural changes of surface
adsorbates at diﬀerent potentials were studied by X-ray
photoelectron spectroscopy.1 The surface distributions ofadsorbates were investigated by scanning tunneling microscopy2 and atomic force microscopy.3 However, the structures of
surface adsorbates under vacuum are diﬀerent from those underatmospheric conditions. There is no information on electrolytes
in bulk phase. It is important to investigate the structures ofelectrolyte molecules at solid−liquid interfaces of electrodes
under realistic conditions.
The structures of solvent water molecules at diﬀerent
potentials were determined from OH stretching mode in
Fourier transform infrared spectroscopy.4,5 The adsorption
structures of water molecules were studied by methods focusing
on solid−liquid interfaces, such as sum frequency generation,6
surface-enhanced Raman,7 and surface-enhanced infrared
spectroscopy.8 The orientation of water molecules at the ﬁrst
layers of Ag(111) electrodes was revealed to be changed at
diﬀerent potentials in 0.1 M NaF by X-ray scattering.9 The
distributions of electrodeposited nanoparticles on electrodesurfaces were observed by transmission electron microscopy.10
The structures of electrolyte molecules at solid−liquidinterfaces of electrodes have been studied extensively by several
methods, but there are few studies by direct measurement of© 2013 American Chemical Society

the change in valence and structure with element-speciﬁc
spectroscopy. It is also diﬃcult to investigate the structure ofelectrolytic solutes in dilute electrolyte solutions.
X-ray absorption spectroscopy (XAS) is a powerful tool tostudy local electronic structures and is applicable to liquids. The
element-speciﬁc spectra of solutes can be observed even in
dilute solutions. Nakai et al. developed a liquid cell for XAS ofelectrolyte solutions and revealed dynamic changes of Mn ionsin solution by measuring Mn K-edge XAS.11 Endo et al. studied
adsorption structures of Br ions on Ag(100) surfaces at
diﬀerent potentials in aqueous NaBr solutions by measuring Br
K-edge XAS.12 The structures of electrodeposited Ni−B ﬁlms
were investigated by Ni K-edge XAS.13 Wu et al. used Cu Kedge XAS to investigate the eﬀect of diﬀerent anions on the
structure of underpotential deposition of Cu on Au.14,15 All of
those measurements were performed in the hard X-ray region,in which XAS spectra are easily measured in transmission or
ﬂuorescence mode due to the high transmittance of hard Xrays. However, XAS in the hard X-ray region is not applicableto electrolytes including chemically important elements such as
C, N, and O. It is necessary to probe these electrolytes during
electrochemical reaction.
The soft X-ray region below ∼2 keV has many chemicallyimportant absorption edges such as C, N, and O K-edges.
Recently, the structure of liquid water has been extensively
studied by O K-edge XAS.16−18 Because the X-ray absorption
Received: May 24, 2013
Revised: July 11, 2013
Published: July 16, 2013
16343

dx.doi.org/10.1021/jp405112r | J. Phys. Chem. C 2013, 117, 16343−16348

The Journal of Physical Chemistry C

Articleprocess occurs within several femto seconds, XAS is able toinvestigate hydrogen-bonding networks in liquid water before
reconstruction of the networks. XAS spectra of liquid samples
can be measured by diﬀerent methods from transmission
mode:19,20 ﬂuorescence yield,17 nonresonant Raman technique,18 total electron yield of liquid microjet,16 and inverse
partial ﬂuorescence yield of liquid microjet.21−24 Comparison
with the transmission spectra is discussed as regards background subtraction, normalization, and saturation-correction.
Note that XAS in bulk liquid phase measured by nonresonant
Raman and inverse partial ﬂuorescence yields is nearly the sameas XAS in transmission mode. Because soft X-rays at the O Kedge are strongly absorbed by water itself in XAS measured intransmission mode, it is necessary to keep the thickness of theliquid layer below a few micrometers.25 Therefore, it is diﬃcult
to measure XAS of liquid samples in transmission mode. It has
not been applied previously to XAS studies of electrochemical
reactions of electrolytes.
Recently, we have developed a liquid cell for XAS intransmission mode.19 The liquid layer is sandwiched between
two Si3N4 membranes and is kept at atmospheric pressure. The
liquid sample is under realistic conditions. The thickness of theliquid layer is controllable between 20 and 2000 nm. We havestudied the hydration structure of diﬀerent cations in aqueous
salt solutions by the O K-edge XAS and the local structure ofmethanol−water binary solutions at diﬀerent mixing ratios by
the C and O K-edge XAS. This liquid cell is also able tomeasure XAS of solid−liquid interfaces by depositing a solid
substrate on the membrane and adjustingt the thickness of theliquid layer as small as possible (<20 nm).
In the present work, we have developed an in situ XAS
measurement system to study electrochemical reactions ofelectrolytes under realistic conditions by using a liquid cell withbuilt-in electrodes. Change in valence of Fe ions in an aqueous
iron sulfate solution at diﬀerent potentials is investigated by Fe
L-edge XAS spectra. The redox reaction of Fe ions is one of themost common electrochemical systems because of theirimportance in a variety of ﬁelds. Previously, Fe redox reactions
were mainly studied by voltammetric methods,26−35 but it isdiﬃcult to know the change in valence of Fe ions at diﬀerent
potentials in dilute electrolyte solutions. The Fe L-edge (700eV) is more sensitive than the Fe K-edge to the 3d valence andspin states of Fe ions.36−43 Here we report direct Fe L-edge
XAS measurement of the change in valence of Fe ions inducedby variation of the potential at a gold electrode.

Figure 1. (a) Schematic of the liquid cell for XAS in transmission
mode. The details of the liquid cell are described in the text. (b)Schematics of three electrodes included in the liquid layer (region III).
The working electrode (WE) is Au-deposited on a Si3N4 membrane.
The counter electrode (CE) is a Pt mesh. The reference electrode(RE) is Ag/AgCl immersed in a saturated KCl solution.

exchanged in situ with a tubing pumping system. The thicknessof the liquid layer is controllable between 20 and 2000 nm bychanging the He pressure in regions II and IV.19 The
absorption of soft X-rays in regions II and IV is small due tothe high transmittance of helium.25 The size of the Si3N4
membrane window between regions I and II is 0.2 mm × 0.2
mm, which is small enough to endure a large diﬀerence inpressure (>1 atm). The soft X-ray beam size on the sample isdetermined by this oriﬁce. Although the thickness of the liquid
layer varies at diﬀerent positions over the 2 mm × 2 mm
membranes, it is possible to measure XAS of the liquid layer at
a constant thickness because of the small beam size.19 Soft Xrays, which pass through region II and the liquid layer (regionIII), are detected by a photodiode (IRD AXUV100) in region
IV. The photon energy resolution at the Fe L-edge was set to
0.7 eV. The XAS spectra are obtained from the measured
transmission signal, by the Lambert−Beer formula ln(I0/I), in
which the current I0 is measured for pure liquid water and thecurrent I is for the liquid sample. The photon energy iscalibrated by the O 1s−π* peak (530.80 eV)45 for O2 gas asmixed in a buﬀer gas with He gas in regions II and IV.
For the investigation of electrochemical reactions, threeelectrodes are included in the liquid layer (region III), as shownin Figure 1b. In the present work, the electrolyte is 0.5 M
aqueous iron sulfate at pH 2.2. The working electrode is a Au

2. EXPERIMENTAL METHODSThe experiments were performed on the soft X-ray undulator
beamline BL3U at the UVSOR-III facility.44 Figure 1a shows
schematics of the present liquid cell. As previously described,19
the liquid cell consists of four regions I, II, III, and IV, separated
by 100 nm-thick Si3N4 membranes (NTT AT Co.). SiC
membranes are used for the N K-edge. Region I is connected tothe beamline under vacuum. Regions II and IV are at
atmospheric pressure of helium buﬀer gas. The ﬂow rate of
the buﬀer gas is changeable by a mass ﬂow controller, and thepressure is adjusted by a needle valve at the gas outlet. The thinliquid layer (region III) is sandwiched between two Si3N4
membranes, each with a window size of 2 mm × 2 mm. The
liquid layer is under atmospheric conditions. Two 100 μm thickspacers are set between the window frames of the membranes,
and the membranes are compressed by sealing o-rings to keep
the thin liquid layer below 2000 nm. Liquid samples can be
16344

dx.doi.org/10.1021/jp405112r | J. Phys. Chem. C 2013, 117, 16343−16348

The Journal of Physical Chemistry C

Articleof aqueous iron sulfate solutions.27,28,33 The peak positions are
summarized in Table 1.

deposit on one of the Si3N4 membranes, which consists of Au(10 nm), Cr (5 nm), and Si3N4 (100 nm) multilayer ﬁlms. This
membrane is one side of the liquid layer and is connected with
the Au tab for electrical conduction. The Teﬂon spacer isplaced on the Si frame, opposite the Au contact. The counterelectrode is a Pt mesh, which is immersed in the sample
electrolyte solution. The reference electrode is Ag/AgCl
immersed in a saturated KCl solution and isolated from theliquid sample by a Teﬂon cover. The potential is controlled byusing a potentiostat (Solartron 1287). We have conﬁrmed fromthe Fe L-edge XAS of the Au-deposited membrane that themembrane has no deposition of Fe metals or compounds after
and also during the Fe redox reaction.

Table 1. Peak Positions of the Fe Redox Reactions in an
Aqueous Iron Sulfate Solution with the Diﬀerent Scanning
Rates Measured by CV and XAS Spectraa
potential (V)Oxidation of Fe(II)0.08 mV/s (XAS)
5 mV/s (CV)20 mV/s (CV)
Reduction of Fe(III)0.08 mV/s (XAS)
5 mV/s (CV)20 mV/s (CV)

3. RESULTS AND DISCUSSION
3.1. Cyclic Voltammetry. Figure 2 shows cyclic
voltammetry (CV) spectra of a 0.5 M aqueous iron sulfate

P1

P2

0.72
0.72
0.70

0.34
0.57

0.29
0.35
0.29

a

Oxidation of Fe(II) ions has two processes. One is the simple
oxidation (P1) and the other involves the sulfate ions (P2). The
reduction of Fe(III) ions is a single process (P1).

3.2. Fe L-edge XAS. Figure 3 shows Fe L-edge XAS spectraof aqueous iron sulfate solutions at diﬀerent potentials. Each

Figure 2. CV spectra of 0.5 M aqueous iron sulfate solutions withscanning rates of (a) 20 and (b) 5 mV/s. The horizontal axis is thepotential of the Au electrode versus the Ag/AgCl reference electrode.
The peak positions of the CV spectra are listed in Table 1.
Figure 3. Fe L-edge XAS spectra of Fe ions in a 0.5 M aqueous ironsulfate at diﬀerent potentials: (a) increasing from 0.00 to 0.90 V and(b) decreasing from 0.90 to −0.40 V.

solution measured by using a conventional electrochemical cell.
The working, counter, and reference electrodes are Au, Pt, and
Ag/AgCl with 3 M NaCl, respectively. The CV spectrum
shown in Figure 2a was measured at a scanning rate of 20 mV/
s. The oxidation of Fe(II) ions to Fe(III) ions with increasing
potential is observed as the peak at 0.70 V. When the potential
is decreased, the reduction of Fe(III) ions to Fe(II) ions isobserved as the peak at 0.29 V.
Figure 2b shows a CV spectrum measured at a scanning rateof 5 mV/s. The reduction of Fe(III) ions to Fe(II) ions occurs
as a single peak at 0.35 V, which is close to that obtained at 20
mV/s. The oxidation of Fe(II) ions to Fe(III) ions shows two
peaks, whose positions are 0.57 and 0.72 V. The peak at 0.72 V
is close to that obtained at 20 mV/s and is attributed to simple
oxidation of Fe(II) to Fe(III). The oxidation process at the low
potential involves the sulfate ions, which aﬀect electrode kinetic
parameters and diﬀusion coeﬃcients34,35 at high concentration

XAS spectrum is measured at a constant potential. The Fe Ledge XAS spectra were measured by increasing the potentialfrom 0.00 to 0.90 V, as shown in Figure 3a. The XAS
measurements continued, while the potential was decreased
from 0.90 to −0.40 V, as shown in Figure 3b. The scanning rateof the potential is roughly estimated to be 0.08 mV/s, which isslower than those of the CV measurements. The XAS L3
spectra have signals from both Fe(II) and Fe(III) ions andshow an isosbestic point, indicating that only two species are
involved.
Figure 4a shows reference Fe L-edge XAS spectra for Fe(II)and Fe(III) ions in aqueous iron sulfate solutions. The XAS
spectrum of Fe(II) ions was measured at no applied potential,
whereas that of Fe(III) ions was obtained at an applied
16345

dx.doi.org/10.1021/jp405112r | J. Phys. Chem. C 2013, 117, 16343−16348

The Journal of Physical Chemistry C

Articleof potential and scanning direction. The uncertainty of themeasured fractions is estimated to be within ±0.02. The
composition is determined from the curve ﬁtting of the XAS
spectra. The amount of Fe(II) ions decreases by oxidation of
Fe(II) to Fe(III) with increasing the potential, which consists of
two processes, as indicated by the change in slope of thecomposition curve at 0.7 and 0.4 V. The formation of Fe(II)from Fe(III) with decreasing potential is linear, indicating asingle-step reduction of Fe(III) ions. The fraction of Fe(II) ions
CFe(II) as a function of potential is ﬁt to a sigmoid proﬁle as
C Fe(II)(E) = Cmin +Cmax − Cmin
1 + exp[α(E − E0)]

(1)

where E0 is the central potential and the terms Cmin and Cmax
are minimum and maximum fractions of Fe(II) ions in the
single process, respectively. The slope of the sigmoid proﬁle isinﬂuenced by term α.
Table 1 shows the central potentials of the Fe redox process
obtained by the ﬁts to sigmoid proﬁles, together with the resultsof the CV spectra at the diﬀerent scanning rates. The oxidationof Fe(II) ions to Fe(III) ions has two processes which occur at
potentials of 0.34 and 0.72 V. The reduction of Fe(III) ions toFe(II) ions occurs at a potential of 0.29 V. Because each XAS
spectrum is obtained at a constant potential, the scanning rateof the potential is quite slow (0.08 mV/s) in the XAS
measurements as compared with the CV results (5 and 20 mV/s). As shown in Table 1, the reduction peak of Fe(III) ionsobtained by XAS is close to that obtained by the CV spectra.
The reduction of Fe(III) ions is a simple process that reaches
an equilibrium even at a scanning rate of 20 mV/s. The proﬁleof the reduction is deviated slightly from the sigmoid proﬁle. It
might be inﬂuenced by the sulfate ions forming somecomplexes with Fe(III) ions.35 Two processes are found in
the case of oxidation of Fe(II) ions. One is a simple oxidation
process of Fe(II) to Fe(III). This peak position is 0.72 V at 0.08
mV/s and is close to that observed in the CV spectra,
suggesting that the simple oxidation process reaches equilibrium. The peak position of the other oxidation process is 0.34
V at 0.08 mV/s. In the CV spectra, the peak position is 0.57 V
at 5 mV/s, and it is not observed at 20 mV/s. The reason why
the peak position varies with the scanning rate is that thisoxidation process does not reach equilibrium. The rate of thisprocess is dominantly inﬂuenced by the sulfate ions, whichaﬀect electrode kinetic parameters and diﬀusion coeﬃcients.27,28,33−35 The present XAS results on the change invalence of Fe ions at diﬀerent scanning rates are consistent with
the CV results.

Figure 4. (a) Reference Fe L-edge XAS spectra for Fe(II) and Fe(III)ions in aqueous iron sulfate solutions. (b) Example of the ﬁtting of the
Fe L-edge XAS spectrum at the potential of 0.45 V in the upward
direction by superposition of the reference spectra.

potential of 0.90 V. As shown in Figure 3a, the amount of
Fe(II) ions decreases to form Fe(III) ions nonlinearly when thepotential is increased. When the potential is decreased, a linearchange of the XAS spectra from Fe(III) ions to Fe(II) ions isobserved, as shown in Figure 3b. To obtain the fraction of
Fe(II) and Fe(III) ions, we ﬁt the Fe L-edge XAS spectra at
diﬀerent potentials to a superposition of the reference spectraof Fe(II) and Fe(III) ions shown in Figure 4a. Figure 4b shows
an example of the least-squares ﬁtting for the Fe L-edge XAS
spectrum measured at a potential of 0.45 V in the upward
scanning direction. There is good agreement with a simple
superposition of the reference spectra.
3.3. Mechanism of Fe Redox Reactions. Figure 5 shows
the fraction of Fe(II) ions (Fe(II) + Fe(III) = 1.0) as a function

4. CONCLUSIONSThe change in valence of Fe ions in aqueous iron sulfate
solutions at diﬀerent potentials under realistic conditions has
been investigated by Fe L-edge XAS in transmission mode.
Each X-ray absorption spectrum was measured at a constantpotential by using a newly developed liquid cell with built-in
electrodes. The Fe redox process is investigated by correlating
the XAS results with the CV spectra. The scanning rate of thepotential used in the XAS study is estimated to be 0.08 mV/s,
which is quite slow compared with that of the CV spectra (5and 20 mV/s). The reduction of Fe(III) ions to Fe(II) ionswith decreasing potential is a single process that reaches
equilibrium even at a scanning rate of 20 mV/s. The oxidationof Fe(II) ions to Fe(III) ions with increasing potential consists

Figure 5. Fraction of Fe(II) ions as a function of potential versus Ag/
AgCl with saturated KCl solutions. The arrows indicate the scanning
direction of applied potential. Each fraction includes an error bar,which is within ±0.02. The oxidation of Fe(II) with increasing
potential has two processes, whereas the reduction of Fe(III) withdecreasing potential is a single process. The central potentials of theseFe redox processes are listed in Table 1.
16346

dx.doi.org/10.1021/jp405112r | J. Phys. Chem. C 2013, 117, 16343−16348

The Journal of Physical Chemistry C

Article(9) Toney, M. F.; Howard, J. N.; Richer, J.; Borges, G. L.; Gordon, J.
G.; Melroy, O. R.; Wiesler, D. G.; Yee, D.; Sorensen, L. B. VoltageDependent Ordering of Water Molecules at an Electrode-Electrolyte
Interface. Nature 1994, 368, 444−446.
(10) De Jonge, N.; Ross, F. M. Electron Microscopy of Specimens in
Liquid. Nat. Nanotechnol. 2011, 6, 695−704.
(11) Nakai, I.; Shiraishi, Y.; Nishikawa, F. Development of a New in
Situ Cell for the X-ray Absorption Fine Structure Analysis of theElectrochemical Reaction in a Rechargeable Battery and Its
Application to the Lithium Battery Material, Li1+yMn2‑yO4. Spectrochim.
Acta, Part B 1999, 54, 143−149.
(12) Endo, O.; Kiguchi, M.; Yokoyama, T.; Ito, M.; Ohta, T. In-Situ
X-ray Absorption Studies of Bromine on the Ag(100) Electrode. J.
Electroanal. Chem. 1999, 473, 19−24.
(13) Bediako, D. K.; Lassalle-Kaiser, B.; Surendranath, Y.; Yano, J.;
Yachandra, V. K.; Nocera, D. G. Structure-Activity Correlations in a
Nickel-Borate Oxygen Evolution Catalyst. J. Am. Chem. Soc. 2012, 134,
6801−6809.
(14) Wu, S.; Lipkowski, J.; Tyliszczak, T.; Hitchcock, A. P. Effect ofAnion Adsorption on Early Stages of Copper Electrocrystallization atAu(111) Surface. Prog. Surf. Sci. 1995, 50, 227−236.
(15) Wu, S.; Shi, Z.; Lipkowski, J.; Hitchcock, A. P.; Tyliszczak, T.
Early Stages of Copper Electrocrystallization: Electrochemical and in
Situ X-ray Absorption Fine Structure Studies of Coadsorption of
Copper and Chloride at the Au(111) Electrode Surface. J. Phys. Chem.
B 1997, 101, 10310−10322.
(16) Smith, J. D.; Cappa, C. D.; Wilson, K. R.; Messer, B. M.; Cohen,
R. C.; Saykally, R. J. Energetics of Hydrogen Bond Network
Rearrangements in Liquid Water. Science 2004, 306, 851−853.
(17) Wernet, P.; Nordlund, D.; Bergmann, U.; Cavalleri, M.; Odelius,M.; Ogasawara, H.; Näslund, L.-Å.; Hirsch, T. K.; Ojamäe, L.; Glatzel,
P.; Pettersson, L. G. M.; Nilsson, A. The Structure of the First
Coordination Shell in Liquid Water. Science 2004, 304, 995−999.
(18) Huang, C.; Wikfeldt, K. T.; Tokushima, T.; Nordlund, D.;
Harada, Y.; Bergmann, U.; Niebuhr, M.; Weiss, T. M.; Horikawa, Y.;
Leetmaa, M.; Ljungberg, M. P.; Takahashi, O.; Lenz, A.; Ojamäe, L.;
Lyubartsev, A. P.; Shin, S.; Pettersson, L. G. M.; Nilsson, A. The
Inhomogeneous Structure of Water at Ambient Conditions. Proc. Natl.
Acad. Sci. U. S. A. 2009, 106, 15214−15218.
(19) Nagasaka, M.; Hatsui, T.; Horigome, T.; Hamamura, Y.; Kosugi,
N. Development of a Liquid Flow Cell to Measure Soft X-ray
Absorption in Transmission Mode: A Test for Liquid Water. J. Electron
Spectrosc. Relat. Phenom. 2010, 177, 130−134.
(20) Schreck, S.; Gavrila, G.; Weniger, C.; Wernet, P. A Sample
Holder for Soft X-ray Absorption Spectroscopy of Liquids inTransmission Mode. Rev. Sci. Instrum. 2011, 82, 103101.
(21) Achkar, A. J.; Regier, T. Z.; Wadati, H.; Kim, Y.-J.; Zhang, H.;
Hawthorn, D. G. Bulk Sensitive X-ray Absorption Spectroscopy Freeof Self-Absorption Effects. Phys. Rev. B 2011, 83, 081106(R).
(22) Gotz, M. D.; Soldatov, M. A.; Lange, K. M.; Engel, N.; Golnak,R.; Könnecke, R.; Atak, K.; Eberhardt, W.; Aziz, E. F. Probing CosterKronig Transitions in Aqueous Fe2+ Solution Using Inverse Partial andPartial Fluorescence Yield at the L-Edge. J. Phys. Chem. Lett. 2012, 3,
1619−1623.
(23) Regier, T. Z.; Achkar, A. J.; Peak, D.; Tse, J. S.; Hawthorn, D. G.
Dark Channel Fluorescence Observations Result from Concentration
Effects Rather than Solvent-Solute Charge Transfer. Nat. Chem. 2012,
4, 765−766.
(24) Soldatov, M. A.; Lange, K. M.; Gotz, M. D.; Engel, N.; Golnak,R.; Kothe, A.; Aziz, E. F. On the Origin of Dips in Total Fluorescence
Yield X-ray Absorption Spectra: Partial and Inverse Partial
Fluorescence Yield at the L-Edge of Cobalt Aqueous Solution.
Chem. Phys. Lett. 2012, 546, 164−167.
(25) Chantler, C. T. Detailed Tabulation of Atomic Form Factors,
Photoelectric Absorption and Scattering Cross Section, and Mass
Attenuation Coefficients in the Vicinity of Absorption Edges in theSoft X-Ray (Z = 30−36, Z = 60−89, E =0.1 keV−10 keV), Addressing
Convergence Issues of Earlier Work. J. Phys. Chem. Ref. Data 2000, 29,
597−1048.

of two processes. The process at a high potential is simple
oxidation of Fe(II) ions that reaches equilibrium even at ascanning rate of 20 mV/s. The oxidation process at a low
potential involves the sulfate ions. The peak positions are
dependent on the scanning rate because the rate of this process
is dominantly inﬂuenced by the sulfate ions, which aﬀect
electrode kinetic parameters and diﬀusion coeﬃcients.27,28,33−35
In the present study, the change in valence of Fe ions at
diﬀerent potentials leads to a drastic change in the XAS spectra.
Therefore, the change in valence of Fe ions can be successfully
studied by measurements of the XAS spectrum at a constantpotential. However with the present method it is diﬃcult toobserve small spectral changes in the XAS spectra at diﬀerent
potentials. For such studies it is necessary to measure changes
in soft X-ray absorption at a single photon energy while
scanning the applied potential. Even a subtle spectral change in
the XAS spectrum can be revealed by measurements of theabsorption diﬀerence with a potential modulation as a functionof photon energy. This will enable us to investigate eﬀects ofthe environment of a target atom and also to determine kinetic
parameters in electrochemical reactions at diﬀerent potential
modulation rates. An apparatus to carry out such detailed
measurements has been developed and will be used in future
experiments.

■

AUTHOR INFORMATION

Corresponding Author*E-mail: nagasaka@ims.ac.jp.
Notes

The authors declare no competing ﬁnancial interest.

■

ACKNOWLEDGMENTS
This work is supported by JSPS Grants-in-Aid for ScientiﬁcResearch (Nos. 23685006 and 23245007). We acknowledge thestaﬀ members of the UVSOR-III facility for their kind support.
A.H. thanks IMS for its support during a research leave inwhich this work was performed.

■

REFERENCES(1) Wakisaka, M.; Suzuki, H.; Mitsui, S.; Uchida, H.; Watanabe, M.
Increased Oxygen Coverage at Pt-Fe Alloy Cathode for the Enhanced
Oxygen Reduction Reaction Studied by EC-XPS. J. Phys. Chem. C
2008, 112, 2750−2755.
(2) Cuesta, A.; Kleinert, M.; Kolb, D. M. The Adsorption of Sulfate
and Phosphate on Au(111) and Au(100) Electrodes: An in Situ STM
Study. Phys. Chem. Chem. Phys. 2000, 2, 5684−5690.
(3) Umeda, K.; Fukui, K. Observation of Redox-State-Dependent
Reversible Local Structural Change of Ferrocenyl-TerminatedMolecular Island by Electrochemical Frequency Modulation AFM.
Langmuir 2010, 26, 9104−9110.
(4) Ashley, K.; Pons, S. Infrared Spectroelectrochemistry. Chem. Rev.
1988, 88, 673−695.
(5) Nakamura, M.; Kato, H.; Hoshi, N. Infrared Spectroscopy of
Water Adsorbed on M(111) (M = Pt, Pd, Rh, Au, Cu) Electrodes in
Sulfuric Acid Solution. J. Phys. Chem. C 2008, 112, 9458−9463.
(6) Nihonyanagi, S.; Ye, S.; Uosaki, K.; Dreesen, L.; Humbert, C.;
Thiry, P.; Peremans, A. Potential-Dependent Structure of theInterfacial Water on the Gold Electrode. Surf. Sci. 2004, 573, 11−16.
(7) Pettinger, B.; Philpott, M. R.; Gordon, J. G., II. Further
Observations of the Surface Enhanced Raman Spectrum of Water on
Silver and Copper Electrodes. Surf. Sci. 1981, 105, 469−474.
(8) Ataka, K.; Yotsuyanagi, T.; Osawa, M. Potential-Dependent
Reorientation of Water Molecules at an Electrode/Electrolyte
Interface Studied by Surface-Enhanced Infrared Absorption Spectroscopy. J. Phys. Chem. 1996, 100, 10664−10672.
16347

dx.doi.org/10.1021/jp405112r | J. Phys. Chem. C 2013, 117, 16343−16348

The Journal of Physical Chemistry C

Article(26) George, P. The Oxidation of Ferrous Perchlorate by Molecular
Oxygen. J. Chem. Soc. 1954, 280, 4349−4359.
(27) Huffman, R. E.; Davidson, N. Kinetics of the Ferrous IronOxygen Reaction in Sulfuric Acid Solution. J. Am. Chem. Soc. 1956, 78,4836−4842.
(28) Beukenkamp, J.; Herrington, K. D. Ion-Exchange Investigationof the Nature of Iron(II) in Sulfuric and Perchloric Acid. J. Am. Chem.
Soc. 1960, 82, 3022−3025.
(29) Wells, C. F.; Salam, M. A. Hydrolysis of Ferrous Ions: a Kinetic
Method for the Determination of the Fe(II) Species. Nature 1965,
205, 690−692.
(30) Bockris, J. O. M.; Mannan, R. J.; Damjanovic, A. Dependence of
the Rate of Electrodic Redox Reactions on the Substrate. J. Chem. Phys.
1968, 48, 1898−1904.
(31) Samec, Z.; Weber, J. The Effect of the Double Layer on the Rateof the Fe3+/Fe2+ Reaction on a Platinum Electrode and theContemporary Electron Transfer Theory. J. Electroanal. Chem. 1977,
77, 163−180.
(32) Weber, J.; Samec, Z.; Mareček, V. The Effect of Anion
Adsorption on the Kinetics of the Fe3+/Fe2+ Reaction on Pt and Au
Electrodes in HClO4. J. Electroanal. Chem. 1978, 89, 271−288.
(33) Iwai, M.; Majima, H.; Awakura, Y. Oxidation of Fe(II) in
Sulfuric Acid Solutions with Dissolved Molecular Oxygen. Metall.
Mater. Trans. B 1982, 13, 311−318.
(34) Gil, A. F.; Salgado, L.; Galicia, L.; González, I. PredominanceZone Diagrams of Fe(III) and Fe(II) Sulfate Complexes in Acidic
Media. Voltammetric and Spectrophotometric Studies. Talanta 1995,
42, 407−414.
(35) Gil, A. F.; Galicia, L.; González, I. Diffusion Coefficients andElectrode Kinetic Parameters of Different Fe(III) - Sulfate Complexes.
J. Electroanal. Chem. 1996, 417, 129−134.
(36) Hitchcock, A. P.; Wen, A. T.; Rühl, E. Transition Metal 2p
Excitation of Organometallic Compounds Studied by Electron Energy
Loss Spectroscopy. Chem. Phys. 1990, 147, 51−63.
(37) George, S. J.; Van Elp, J.; Chen, J.; Ma, Y.; Chen, C. T.; Park, J.B.; Adams, M. W. W.; Searle, B. G.; De Groot, F. M. F.; Fuggle, J. C.;
Cramer, S. P. L-Edge X-ray Absorption Spectroscopy of Pyrococcus
furiosus Rubredoxin. J. Am. Chem. Soc. 1992, 114, 4426−4427.
(38) Van der Laan, G.; Kirkman, I. W. The 2p Absorption Spectra of
3d Transition Metal Compounds in Tetrahedral and Octahedral
Symmetry. J. Phys., Condens. Matter 1992, 4, 4189−4204.
(39) Wen, A. T.; Rühl, E.; Hitchcock, A. P. Inner-Shell Excitation ofOrganoiron Compounds by Electron Impact. Organometallics 1992,
11, 2559−2569.
(40) Cressey, G.; Henderson, C. M. B.; Van der Laan, G. Use of Ledge X-ray Absorption Spectroscopy to Characterize Multiple ValenceStates of 3d Transition Metals; a New Probe for Mineralogical andGeochemical Research. Phys. Chem. Miner. 1993, 20, 111−119.
(41) De Groot, F. M. F. X-ray Absorption and Dichroism ofTransition Metals and Their Compounds. J. Electron Spectrosc. Relat.
Phenom. 1994, 67, 529−622.
(42) Huse, N.; Kim, T. K.; Jamula, L.; McCusker, J. K.; De Groot, F.
M. F.; Schoenlein, R. W. Photo-Induced Spin-State Conversion in
Solvated Transition Metal Complexes Probed via Time-Resolved Soft
X-ray Spectroscopy. J. Am. Chem. Soc. 2010, 132, 6809−6816.
(43) Lange, K. M.; Kothe, A.; Aziz, E. F. Chemistry in Solution:
Recent Techniques and Applications Using Soft X-ray Spectroscopy.
Phys. Chem. Chem. Phys. 2012, 14, 5331−5338.
(44) Hatsui, T.; Shigemasa, E.; Kosugi, N. Design of a Transmission
Grating Spectrometer and an Undulator Beamline for Soft X-ray
Emission Studies. AIP Conf. Proc. 2004, 705, 921−924.
(45) Coreno, M.; de Simone, M.; Prince, K. C.; Richter, R.;
Vondrácě k, M.; Avaldi, L.; Camilloni, R. Vibrationally Resolved
Oxygen K → Π* Spectra of O2 and CO. Chem. Phys. Lett. 1999, 306,
269−274.

16348

dx.doi.org/10.1021/jp405112r | J. Phys. Chem. C 2013, 117, 16343−16348