New Insights into Pseudocapacitive Charge-Storage Mechanisms in Li

Article pubs.acs.org/JPCC

New Insights into Pseudocapacitive Charge-Storage Mechanisms in Li-Birnessite Type MnO2 Monitored by Fast Quartz Crystal Microbalance Methods Carlota Ridruejo Arias,†,‡ Catherine Debiemme-Chouvy,†,‡ Claude Gabrielli,†,‡ Christel Laberty-Robert,§ Alain Pailleret,†,‡ Hubert Perrot,†,‡ and Ozlem Sel*,†,‡ †

Sorbonne Universités UPMC Univ Paris 06, UMR 8235, Laboratoire Interfaces et Systèmes Electrochimiques, F-75005, Paris, France ‡ CNRS, UMR 8235, LISE, F-75005, Paris, France § LCMCP-CNRS-UMR-7574-Collège de France, Bat. C-D, 11, place Marcelin Berthelot, 75231 Paris, France S Supporting Information *

ABSTRACT: Fast ionic transfer and transport properties continue to be one of the main pressing research concerns regarding energy storage materials (batteries, supercapacitors). Accompanying this search for optimal materials, appropriate characterization tools to assess key parameters of newly developed materials are required. In spite of its great relevance, fast electrogravimetric methods, i.e., coupling fast quartz crystal microbalance (QCM) and electrochemical impedance spectroscopy (EIS) (ac-electrogravimetry), have not yet been used for studying transfer and transport phenomena in materials for charge storage (except for the use of QCM along with cyclic voltammetry experiments (EQCM)). This coupled method, socalled ac-electrogravimetry, differs from classical EQCM and measures the usual electrochemical impedance, ΔE/ΔI (ω), and the mass variations of the film under a sinusoidal potential perturbation, Δm/ΔE (ω), simultaneously. This coupling has the ability to detect the contribution of the charged or uncharged species and to separate the anionic, cationic, and free solvent contributions during the various (pseudo)capacitive processes. The Li-birnessite type MnO2 thin films were studied in two different media, LiClO4 and NaClO4, by ac-electrogravimetry. The Li+ ions, Na+ ions, and their respective hydrated ionic species are detected to be involved in the pseudocapacitive charge storage of Li-birnessite type MnO2 thin films. The kinetics (fci) and resistance (Rti) of charged and noncharged species transferred at the electrode/electrolyte interfaces and the number of water molecules in the hydration shell of the ions are estimated considering integer values. The opposite flux direction of free water molecules was also detected by ac-electrogravimetry. This indicates that there is a population of hydrated Li+ or hydrated Na+ ions losing their hydration shell before being transferred at the electrode/electrolyte interfaces. Therefore, the effect of desolvation is clearly and experimentally demonstrated. The ac-electrogravimetry responses of the electrodeposited Li-birnessite type MnO2 thin films can serve as a gravimetric probe for studying the charge-storage mechanisms and extracting subtleties unreachable with classical tools.

■

(MnO2 )bulk + C+ + e− ↔ MnOOC

INTRODUCTION Layered transition-metal oxides such as birnessite type manganese oxides are very appealing for charge storage1,2 because of their ability to intercalate ions in a wide range of sites.3,4 Both faradaic and nonfaradaic mechanisms can participate to the charge storage. Two different charge-storage routes are feasible when a material is polarized in an electrolyte: (i) cations electrochemically adsorbed onto the surface (nonfaradaic) or (ii) cations intercalated into the interlayer gaps of a layered material (faradaic). These two charge-storage mechanisms in MnO2 can be described by reactions I and II:5 +

−

− +

(MnO2 )surface + C + e ↔ (MnO2 C )surface © 2014 American Chemical Society

(II)

The faradaic charge-transfer mechanism (II) was also classified according to the site where it takes place, either on the surface or in the interlayer gaps, and described as (i) redox pseudocapacitance or (ii) intercalation pseudocapacitance, respectively. From a chemical perspective, reactions (i) and (ii) are very similar in that the electrochemical ion (C+: Li+, Na+, K+, H+...) faradaic contribution is associated with the Received: August 23, 2014 Revised: October 23, 2014 Published: October 23, 2014

(I) 26551

dx.doi.org/10.1021/jp508543h | J. Phys. Chem. C 2014, 118, 26551−26559

The Journal of Physical Chemistry C

Article

(QCM) and electrochemical impedance spectroscopy (EIS) (ac-electrogravimetry), have not yet been used for studying transfer and transport phenomena in materials for charge storage (except for the use of QCM along with cyclic voltammetry experiments (EQCM)8,9,11−13). Here we propose an alternative characterization tool to overcome the limitations of the classical EQCM for studying charge-storage mechanisms. The pseudocapacitive behavior in Li-birnessite type layered MnO2 thin-film electrodes was investigated by coupled characterization methods resolved at a temporal level (fast QCM−electrochemical impedance spectroscopy). This coupled method, ac-electrogravimetry, simultaneously measures the electrochemical impedance, ΔE/ΔI (ω), and the mass/ potential transfer function (Δm/ΔE (ω)) during a sinusoidal potential perturbation with a small amplitude applied to the modified electrode.14−16 This coupling has the ability to detect the contribution of the charged or uncharged species and to separate the anionic, cationic, and free solvent contributions during the various (pseudo)capacitive processes. The Libirnessite type MnO2 thin films were studied in two different media, LiClO4 and NaClO4, to understand the charge-storage mechanisms and to identify the species contributions to the charge compensation process. It has been demonstrated herein that the ac-electrogravimetry responses of the electrodeposited Li-birnessite type MnO2 thin films can serve as a gravimetric probe for studying the complex charge-storage mechanisms and extracting subtleties unreachable with classical tools.

reduction of either metal atoms located on the surface or in the interlayer lattice planes.6 The faradaic and nonfaradaic reactions were broadly investigated in the past decades. The intercalation of alkali metal ions into birnessite type manganese oxides in aqueous media studied by cyclic voltammetry and chemical composition analysis suggested that protons but not the alkali metal cations are directly intercalated. The presence of the alkali metal cations in the oxide matrix was explained by an ion exchange reaction between the protons and the cations.7 The X-ray photoelectron spectroscopy (XPS) study by Brousse et al. on reduced manganese oxide in aqueous Na2SO4 electrolyte indicated that the Na/Mn ratio was much lower than what was anticipated for charge compensation dominated by Na+. Therefore, the involvement of protons in the charge compensation was suggested to explain the low Na/Mn ratio in the reduced electrodes. Wu and co-worker investigated the pseudocapacitive charge-storage reaction of MnO2·nH2O in aqueous alkali and alkaline salt solutions by EQCM and related the results of mass variations (Δm) to XPS and in situ synchrotron X-ray diffraction (XRD) results.8 The resulting mass changes, Δm, of the MnO2·nH2O films were monitored during cyclic voltammetry measurements (EQCM) and massto-charge ratios (MCRs) were calculated in various aqueous media such as LiCl, NaCl, KCl, CsCl, and CaCl2 electrolytes. If the pseudocapacitance is attributed to either electroadsorption of alkali cation or protonation reaction balancing valence variation at the Mn sites, the former is expected to give a MCR equal to the equivalent weight of the alkali ion, while the latter should give a MCR of 1 g mol−1. However, none of the measured MCRs in the work of Wu and co-worker matches the equivalent weight of either H+ or alkali ions (Li+, Na+, K+, Cs+ or Ca2+).8 The measured MCR from EQCM was ∼19 g mol−1 in aqueous NaCl electrolytes. The XPS study also did not detect any Na in the reduced electrodes, although a low Na/ Mn ratio was detected by Brousse and co-workers.5 Therefore, hydronium ions (H3O+) were suggested to intervene almost exclusively in the pseudocapacitive charge-storage mechanism. The MCR values higher than the equivalent weight of H3O+ were treated using a simple calculation based on the combination of two ions, hydronium ions and corresponding alkali cations. MCR values were ∼16 g mol−1 in LiCl electrolytes, but in that case XPS study could not be used to support the EQCM because XPS atomic sensitivity factor for Li (1 s) is very small, so data cannot give unambiguous evidence of presence or absence of Li in the reduced film electrode. The authors stated that charge transfer at the Mn sites upon reduction/oxidation is balanced by insertion/extraction of the solution cations into/from the oxide structure but H3O+ plays a predominant role in all cases.7,8 The detailed study of Wu and co-worker is one of the first attempts to explain pseudocapacitive charge-storage reaction by EQCM, specifically for MnO2· nH2O films. However, the limitations of the technique appear here clearly. EQCM gives a global response corresponding in fact to several possible pathways such that ions, ions with solvation shells, and free solvent molecules can contribute to the electrochemical process. Additionally, ions could lose a part of their solvation to access to the sites in smaller nanopores such as described for the supercapacitive behavior of porous carbide-derived carbons (CDC) in micropores by Simon and co-workers.9,10 In spite of its great relevance, fast electrogravimetric methods, i.e., coupling fast quartz crystal microbalance

■

EXPERIMENTAL SECTION Thin Film Synthesis. The gold-patterned quartz substrates of 9 MHz (Temex, France) were used as working electrodes. A platinum grid and an Ag/AgCl (3 M KCl saturated with AgCl) served as counter electrode and reference electrode, respectively. Li-birnessite type MnO2 films were prepared by potentiostatic electrodeposition at 1.0 V vs Ag/AgCl in a bath containing 2 mM MnSO4 (AnalaR NORMAPUR, ≥99%) and 50 mM LiClO4 (Aldrich, 99.99% trace metals basis) as described previously in the literature.17,18 During electrodeposition, a salt-bridge junction equipped with a porous glass frit on the end was used to prevent the reference electrode from being contaminated by the media and vice versa. The electrodeposition was performed by using a potentiostat (Autolab PGSTAT100), and the deposition process was monitored by QCM measurements. The type of intercalated cation in the birnessite structure is dictated by the composition of the electrolyte. Film area is 0.25 cm2, and the film thickness was controlled by the electrodeposition time, typically 30 min. The electrochemical synthesis was followed by a mild thermal treatment at 80 °C for 12h. Structural and Electrochemical Characterization. The film morphology and the thickness were investigated by field emission gun scanning electron microscopy (SEM−FEG) (Zeiss, Supra 55). The synthesis was monitored by a labmade QCM device providing the electrodeposited mass to be determined by the frequency variation of the quartz crystal resonator based on the Sauerbrey equation.19 Crystallinity was studied in ambient conditions by X-ray diffraction (Phillips PANanalytical X’Pert Pro) using Cu Kα radiation (λ = 1.54184 Å). The oxidation state of the manganese was determined by Xray photoelectron spectroscopy (on a VG ESCALAB 250i-XL spectrometer using monochromatic Al Kα radiation as the Xray source). The electrochemical experiments were carried out in a classical three-electrode cell using aqueous solutions of 0.5 26552



Article

Figure 1. Working electrode coated with Li-birnessite type MnO2: (a) FEG-SEM image of the working electrode (blue arrow is the gold electrode of the quartz resonator and red arrow shows the electroactive thin film) and (b) schematic representation of the electrochemical measurement principle and the structure of Li-birnessite type MnO2 stacked along the c axis demonstrating the layered arrangement of oxygen (red) and manganese (inside octahedra) edge-sharing MnO6 octahedra intercalated with cations. k2

M LiClO4 and NaClO4 as electrolytes. The reference electrode was Ag/AgCl (3 M KCl saturated with AgCl) electrode; the counter electrode was a platinum grid, and the working electrode was the modified gold-patterned quartz. During electrochemical tests, a salt-bridge junction equipped with a porous glass frit on the end was used to prevent the reference electrode from being contaminated by the media and vice versa. For ac-electrogravimetry, a four-channel frequency response analyzer (FRA, Solartron 1254) and a potentiostat (SOTELEM-PGSTAT) were used. The QCM was used under a dynamic regime; the modified working electrode (0.2 cm2) was polarized at a selected potential, and a sinusoidal small amplitude potential perturbation (50 mV rms) was superimposed. The microbalance frequency change, Δf m, corresponding to the mass response, Δm, of the modified working electrode was measured simultaneously with the ac response, ΔI, of the electrochemical system. The resulting signals were sent to the four-channel FRA, which allowed the electrogravimetric transfer function, (Δm/ΔE)(ω), and the electrochemical impedance, (ΔE/ΔI)(ω), to be simultaneously obtained at a given potential.

⟨P2⟩ + e− + c2+ ⇄ ⟨P2c2⟩ k 2′

where ⟨P1c1⟩ and ⟨P2c2⟩ are the film matrices doped with cations. The cation transfers at the film|solution interface are taken into account only as rate-limiting steps because the cation transports inside the thin film and in the solution are supposed to be fast enough through thin films or in sufficiently concentrated electrolytes. Under the effect of a sinusoidal potential perturbation with low amplitude, ΔE, imposed on the electrode−film−electrolyte system, sinusoidal fluctuations of concentration, ΔCi, are observed. In the present case, for electroactive metal oxide thin films, the change of the concentration, ΔCi, of each species (cation1 (c1), cation 2 (c2), and free solvent (s)) with potential ΔE can be calculated using eq 3

■

where df is the film thickness and Ki and Gi are the partial derivatives of the flux, Ji, eqs S5−S7 in Supporting Information, with respect to the concentration and the potential, respectively, Ki = (∂Ji/∂Ci)E and Gi = (∂Ji/∂E)Ci, where Ji stands for the flux of the species i crossing the film/electrolyte interface. More precisely, Ki is the kinetics rate of transfer and Gi is the inverse of the transfer resistance, Rti, of the species at the film/electrolyte interface (where i is the cation c1 or c2 or the free solvent s). The charge/potential transfer function, (Δq/ΔE)|th, is calculated for the insertion/expulsion of the two cations, c1 and c2, using the Faraday number, F, and the film thickness, df

THEORETICAL BACKGROUND The transfer of two cations, c1 and c2, in the electroactive films incorporating two different sites, P1 and P2, during the redox reaction of the host material ⟨P⟩ where a single electronic transfer takes place in two independent sites can be expressed as k1

⟨P1⟩ + e− + c1+ ⇄ ⟨P1c1⟩ k1′

(2)

(1)

and 26553



Article

Figure 2. Electrochemical characterization of Li-birnessite type MnO2 in aqueous electrolytes (0.5 M LiClO4 and NaClO4): (a) cyclic voltammetry curves measured at a scan rate of 25 mV s−1 and simultaneously monitored microbalance mass variations, (b) specific capacitance values (calculated from panel a) as a function of selected scan rate, and (c and d) Fdm/dQ values as a function of the potential obtained from the reduction and oxidation branch of the EQCM results, respectively.

Δq ΔE

⎛ ⎞ Gc1 Gc2 − (ω) = −Fdf ⎜ − ⎟ jωdf + Kc2 ⎠ ⎝ jωdf + Kc1 th

(4)

The second main transfer function, (Δm/ΔE)|th, can be calculated theoretically taking into account the solvent contribution through the two parameters Ks and Gs: Δm ΔE

■

⎛ Gc1 Gc2 + mc2 (ω) = df ⎜mc1 jωdf + Kc2 ⎝ jωdf + Kc1 th

+ ms

⎞ Gs ⎟ jωdf + K s ⎠

c1s

Δm ΔE

c1s

(ω) = th

th

Δm ΔE

(5)

(ω) + th

mc2 Δq F ΔE

(ω) th

⎛ ΔC ΔC ⎞ (ω) = df ⎜(mc1 − mc2) c1 + ms s ⎟ ⎝ ΔE ΔE ⎠

c2s

Δm ΔE

c2s

(ω) = th

th

Δm ΔE

(ω) + th

mc1 Δq F ΔE

(ω) th

⎛ ΔC ΔC ⎞ (ω) = df ⎜(mc2 − mc1) c2 + ms s ⎟ ⎝ ΔE ΔE ⎠

(8)

(9)

RESULTS AND DISCUSSION Synthesis and Structural Characterization. Thin films of Li-birnessite type MnO2 in our study (Figure 1a) are synthesized on the gold electrode of a quartz resonator in a one-step electrodeposition method from solutions containing manganese and lithium salts (see Experimental Section and refs 17 and 18 for details). Figure S1a in Supporting Information shows typical current and microbalance frequency profiles, simultaneously measured on a quartz crystal during electrodeposition. The total mass of the film is typically ∼100 μg under the electrodeposition conditions described here. This is determined by converting the frequency variations of the quartz crystal to the mass variations by using the Sauerbrey equation19 (Figure S1b in Supporting Information). The weight versus charge density curve is calculated from Figure S1a in Supporting Information allows the F(dm/dQ) function to be determined, which is approximately equivalent to the molecular mass of the deposited film (Figure S2 in Supporting Information). Mexperimental = 96.1 g mol−1 is higher than the w MnO theoretical molecular mass of the Mw 2 = 86.9 g mol−1, which indirectly shows the incorporation of Li+ ions and water

where mc1, mc2, and ms are the atomic weight of each involved species. Partial mass/potential TF are also estimated either by removing the c2 contribution, calculating (Δm/ΔE)|c1s th or the c1 contribution, calculating (Δm/ΔE)|c2s th . In the case of cation c1 and the solvent, the following equations are used: Δm ΔE

Δm ΔE

(6)

(7)

or for cation c2 and the free solvent 26554



Article

Figure 3. Two main transfer functions, (a) (Δq/ΔE)(ω) and (b) (Δm/ΔE)(ω), and two partial transfer functions, (c) (Δm/ΔE)|c1s(ω) and (d) (Δm/ΔE)|c2s(ω), involving three species, specifically Na+, Na+·H2O, and H2O (both experimental and theoretical curves are given); (e) schematic representation of the ion insertion process. Measurements are in 0.5 M NaClO4 electrolyte at 0.6 V vs Ag/AgCl.

Information was used to determine the mean oxidation state of the manganese, which was estimated as ∼3.8 (see Figure S5 in Supporting Information). When the XPS results and the IV molecular weight of LixMnIII x Mn1−xO2·nH2O estimated from experimental = 96.1 g mol−1) are combined, the EQCM study (Mw IV the composition of the film should be Li0.2MnIII 0.2Mn0.8O2·nH2O, and the value of n is ∼0.5 (see eq S2 in Supporting Information) which is in agreement with previous reports. Electrochemical Characterization. Electrochemical Quartz Crystal Microbalance (EQCM). Figure 2a shows the cyclic voltammetry (CV) curves of Li-birnessite type MnO2, indicating a classical pseudocapacitive behavior with a quasirectangular shape. However, the electrochemical behavior in LiClO4 exhibits slight differences compared to that in the NaClO4 media. The current values measured are two times higher and a faradic contribution is more clearly observed at the potential values around 0.4 V vs Ag/AgCl. The capacitance values (Figure 2b) are gravimetrically normalized because of the in situ mass estimation of the film (QCM measurements during electrodeposition, Figure S1b in Supporting Informa-

molecules into the manganese oxide during electrodeposition. The FEG-SEM image of the film (Figure 1a and Figure S3a in Supporting Information) is representative of a sheetlike structure (Figure 1b) which is consistent with previous reports. The average film thickness is ∼400 nm, and it can be tailored by changing the electrodeposition time. The EDX analysis indicates the presence of the Mn and O in the birnessite type MnO2 thin films (Figure S3b in Supporting Information). The other elements present in the film are not detected because of their small quantities and/or the detection limit of the technique. The XRD pattern of the electrodeposited film (Figure S4 in Supporting Information) reveals distinct peaks at 2θ = 12.2°and 2θ = 24.5° which can be assigned to the (001) and (002) planes of crystalline birnessite type layered structure.17,18 The XPS survey spectrum and O 1s core level spectra of Li-birnessite type MnO2 are presented in Figure S5 in Supporting Information. The O 1s core level spectra were used to estimate the mean oxidation state of manganese in Libirnessite type MnO2 similar to the procedure described by Brousse and co-workers. 5 Equation S1 in Supporting 26555



Article

Figure 4. Evolution of the characteristic frequency, f i, and of the transfer resistance, Rti, for various species in two different aqueous electrolytes, (a, c) 0.5 M LiClO4 and (b, d) 0.5 M NaClO4, at different applied potentials.

scenarios actually takes place. Therefore, an ac-electrogravimetry study (shown schematically in Figure S6 in Supporting Information section D) is performed to determine the insertion−expulsion mechanisms of the pseudocapacitive Libirnessite type MnO2 thin films. AC-Electrogravimetry Study. In NaClO4 Solutions. Experiments were given in priority in NaClO4 media as the electrolyte appears to be an attractive candidate for supercapacitor applications. Figure 3 depicts an example of the transfer functions (TFs) obtained from ac-electrogravimetry in 0.5 M NaClO4 electrolyte at 0.6 V vs Ag/AgCl. The charge/potential TF ((Δq/ΔE)(ω)), Figure 3a) presents mainly two loops in the experimental curve indicating that two charged species intervene (see Theoretical Background). In the mass/potential TF ((Δm/ΔE)(ω), Figure 3b), one big loop appears in the third quadrant at high and medium frequencies (HF and MF). This loop can be attributed to either one species or to two species whose time constants are not sufficiently different from each other. The loops in the third quadrant are characteristic of cation contributions or free solvent molecules in the same flux direction. Another contribution also appears at lower frequencies (LF) in the fourth quadrant (either anions or water molecules with opposite flux direction), highlighting the challenge in obtaining an exact identification of these two or three loops (Figure 3b). Therefore, several configurations were tested using theoretical functions (eqs 4, 5, 7, and 9) to determine the exact contribution of each species. Because the electrolyte is aqueous 0.5 M NaClO4, the first hypothesis is to assume that only Na+ ions and free solvent molecules are involved (directly or indirectly) in the charge compensation. It can be seen in Figure S7 in Supporting Information that the theoretical TF (obtained using eq 4) based only on one cation contribution cannot fit the experimental (Δq/ΔE)(ω) TF: one loop is missing and the

tion). The classical decrease of the specific capacitance as a function of the scan rate is observed. In LiClO4 media, the capacitance values are higher than that measured in NaClO4, which is in good agreement with the higher current values observed in aqueous LiClO4 solutions (Figure 2a). The ions or hydrated ions are inserted or expelled and can be accompanied by free solvent molecules which can contribute to the QCM responses in Figure 2a. The QCM response is amplified (∼4 times) when cycled in aqueous LiClO4 solutions. Normally, the differences in the current values and the different molecular weight values between Li+ and Na+ ions should counterbalance and practically should result in similar mass variations. The significantly higher mass variations in LiClO4 media may indicate the contribution of the free solvent molecules or higher hydration level of Li+ ions compared to Na+ ions when cycled in LiClO4 and NaClO4 media, respectively. This result indicates a mechanism of ion transfer more complex than we can imagine in an approach often described in the literature (eqs I and II). The F(dm/dQ) values as a function of the potential are obtained from the reduction and the oxidation branch of the EQCM and are presented in Figure 2c,d (details of calculations are given in the Supporting Information). As indicated by the higher mass uptake of the Li-birnessite type MnO2 thin films in aqueous LiClO4 media, the absolute values of F(dm/dQ) values are higher, especially for the oxidation branch (Figure 2c,d). If there is strictly one species involved, this function should be equivalent to the molecular mass of this species inserted or expelled. However, for a complex system, like that observed in pseudocapacitive materials, it corresponds to an average molecular weight related to the various species in terms of mass and kinetics. These results are somewhat in line with the previous reports of Wu and co-worker8 but clearly show the limitations of the classical EQCM technique. The EQCM does not provide unambiguous information on which of the possible 26556



Article

correspond to the water molecules caused by further desolvation of the hydrated Na+ ions transferred at the electrode/ electrolyte interface. In that way, these ions can access the sites in the film smaller than the size of the hydrated Na+ ions following route A presented in Figure 3e. Alternatively, these expelled water molecules can be those already present in the interlayer distances of the film. They can be expelled from the film for steric hindrance reasons to liberate the space during charge compensation, as shown schematically in Figure 3e, route B. Our experimental data show that the transfer resistances of these water molecules (Figure 4) are significantly higher than the transfer resistance which would correspond to the exclusion of free water molecules. Therefore, situation A in Figure 3e is more likely (further desolvation of Na+·nH2O). It is important to note that the number of molecules of water hydrating Na+ ions are n = 1 at 0.6 V vs Ag/AgCl. The values for bulk hydration numbers of Na+ ions vary in the literature between four and eight molecules.20 This indicates that the Na+ ions loose a part of their hydration shell before being transferred at the electrode/electrolyte interfaces (steps I and II in Figure 3e). The further dehydration may occur once inside the film, which corresponds to the step III situation A in Figure 3e. To our knowledge, this is the first time that a fair and in situ determination of this effect of desolvation has been clearly demonstrated. Figure 3 constitutes an example of the experimental and theoretical TFs obtained from ac-electrogravimetry study of Libirnessite type MnO2 films measured in 0.5 M NaClO4 electrolyte at 0.6 V vs Ag/AgCl. These measurements were performed as a function of the applied potential in 0.5 M NaClO4 and LiClO4 aqueous solutions. The f i and Rti of the transfer of various sodium- and lithium-containing species are presented in Figure 4. The f i and Rti of the transfer of sodium species (Figure 4b,d) do not significantly vary as a function of applied potential. This agrees well with the rectangular shape of the CV profile shown in Figure 2a, indicating the presence of electroadsorption rather more significant than faradaic reaction. The kinetic parameters and the transfer resistance values independent of the applied potential may indicate that the principle contribution in the capacitive charge-storage mechanism is based on electroadsorption (reaction mechanism II (Introduction)) of the sodium and hydrated sodium ions (Figure 4). This is also supported by the capacitance values calculated from acelectrogravimetry data for hydrated and nonhydrated Na+ ions, which are fairly constant as a function of applied potential (Table S2 in Supporting Information). Faradaic contribution can be neglected here as the redox reactions are correlated to Li+ ions (already inserted in the MnO2 lattice during the synthesis), and probably they are not correlated to Na+ ions which are poorly size compatible. However, hydrated Na+ ions are the faster species at all potentials studied (Figure 4b), and their transfer resistance is smaller than that of Na+ ions alone (Figure 4d). At first sight, this result can be considered to be in stark contradiction with the fact that the smaller ions are transferred more rapidly. The hydrated sodium ions are electroadsorbed easily on the accessible surface sites; thus, their transfer is relatively facile and it is not necessary to completely remove their hydration shell (Figure 5a). In contrast, the nonhydrated sodium ions must access the internal sites to be electroadsorbed, but for this process they must completely lose their hydration shell (Figure

frequencies do not correspond to each other. Therefore, this mechanism is excluded. It is evident that more than one cation contribution is required for the charge compensation. The second hypothesis considers two cation contributions, specifically Na+ and hydrated Na+ ions (Na+·nH2O with n = 1 in this simulation) without free solvent contribution. Figure S8A in Supporting Information shows a good agreement between the experimental and theoretical charge/potential TFs ((Δq/ΔE)(ω)) (obtained using eq 4). The mass/potential TF ((Δm/ΔE)(ω)) given in Figure S8B in Supporting Information shows that the theoretical (obtained using eq 5) and experimental data correspond well at the HF region, but the LF region is not fitted with this hypothesis. Similarly, partial mass/potential TFs do not agree with the theoretical curves (Figure S8C,D in Supporting Information). Various configurations are tested and eliminated, which leads us to the theoretical TFs shown in Figure 3 based on three species. In this way, ambiguities concerning the various different contributions can be excluded, and this methodology was used for all the other applied potentials. In this consideration, the theoretical TFs involve both Na+ ions and hydrated Na+ ions for the charge compensation in pseudocapacitive reactions. The H2O molecules (free solvent molecules) indirectly intervene. On the basis of this hypothesis, the experimental and theoretical (Δq/ΔE)(ω) TFs (obtained using eq 4) (Figure 3a) are in good agreement in terms of frequencies: two well-separated loops correspond to the hydrated Na+ and Na+ ions alone at the HF and LF region, respectively. The experimental (Δm/ΔE)(ω) transfer function (Figure 3b) also agrees well with the theoretical (Δm/ΔE)(ω) TF (obtained using eq 5). The small contribution in the fourth quadrant is related to the free solvent molecules, which have an opposite flux with the cationic species (Na+ ions and hydrated Na+ ions) under these experimental conditions. To validate our hypothesis involving three different species, partial mass/ potential TFs were also analyzed. Partial mass/potential TFs are estimated, for example, by removing the c2 contribution and calculating (Δm/ΔE)| thc1s or by removing the c1 contribution and calculating (Δm/ΔE)|c2s th (eqs 7 and 9). Figure 3c,d exhibits a good agreement between the theoretical and experimental data. These partial mass/potential TFs provide a crosscheck for validating the hypothesis involving three different species and a better separation of the various contributions. The characteristics of their transfers, in terms of parameter values, are presented in Table S1 in Supporting Information. The Ki and Gi constants are determined by fitting the experimental data using the theoretical functions for (Δq/ ΔE)(ω) and (Δm/ΔE)(ω)TFs (eq 4 and 5) The equivalent weight of the charged and uncharged species are determined by (Δm/ΔE)(ω), which provide the identification of the species (Table S1 in Supporting Information). The K i and G i parameters permit the calculation of the characteristic frequency ( f i) and the transfer resistance (Rti) of the species described in eqs S12 and S13 in Supporting Information, respectively. The characteristic frequencies at 0.6 V vs Ag/AgCl (fc1 and fc2, corresponding to the transfer of Na+ ions and hydrated Na+ ions) are given in Figure 4, indicating that under the present experimental conditions hydrated Na+ ions are transferred faster than Na+ ions (fc2 > fc1). Accordingly, the Rti of the transfer of hydrated Na+ ions is significantly lower, indicating the ease of their transfer compared to that of Na+ ions alone. The water molecules have flux directions opposite to that of the cationic species. These water molecules may 26557



Article

Figure 5. Schematic representation of the sodium ion species transferred at the electrode/electrolyte interfaces. Two different species detected by ac-electrogravimetry, specifically Na+ and hydrated Na+ ion species, are shown. Na+·(n − x)H2O have faster kinetics than Na+ probably because it does not completely lose its hydration shell (a). Na+ has slower kinetics because it loses its hydration shell completely before being transferred on the internal sites of the material (b).

Information). For the anodic potential region (0.6 and 0.8 V vs Ag/AgCl) when Li+ species are inserted, water molecules are expelled from the film as in the case of Na+ species. However, for the cathodic potential region (0.2 and 0.4 V vs Ag/AgCl), when Li+ species are inserted, water molecules are also inserted in the film. The kinetics of the water transfer is very close to those of Li+ ions according to the values given in Figure 4a. This indicates that small Li+ ions are accompanied by free water molecules to occupy the free interlayer space liberated by water expulsion occurring at more anodic potentials (0.6 and 0.8 V vs Ag/AgCl) to maintain the lamellar structure of the birnessite layer. This was not the case in NaClO4 media probably because of the larger ion size of Na+. Compared to the hydrated Na+ species, hydrated Li+ species are transferred with more of their hydration shell, which is in agreement with their corresponding hydration energy values.21 In other words, Li+ ion species are more tightly bonded to their water molecules in their hydration shell, as shown experimentally in mi values in Table S1 in Supporting Information.

5b), which will slow down their kinetics and decrease the ease of transfer (Figure 4b,d). The water molecules have the opposite flux direction with sodium species at all potentials. Contrary to Rti values of sodium species, Rti values of water depend strongly on the applied potential (Figure 4d). These changes can be related to the strong interactions between water molecules in a confined medium at more anodic potentials (0.8 and 0.6 V vs Ag:AgCl in Figure 4d). For the ac-electrogravimetry measurements at these potentials, the interlayer spaces are filled with water. However, in the cathodic range, the number of water molecules present in the film is already decreased because of the expulsion process. This results in weaker interactions between water molecules, leading to their relatively easier transfer (lower transfer resistance, Figure 4d). In LiClO4 Solutions. Two different cations, specifically Li+ and hydrated Li+ ions, were also determined by ac-electrogravimetry for Li-birnessite type MnO2 thin films in 0.5 M LiClO4 aqueous solutions. Because of the high hydration energy value of Li+, corresponding dehydration is more difficult and the water molecules are strongly associated. Therefore, the kinetics of Li+ transfer should be slower and their transfer resistance should be higher compared to those of Li+·H2O. However, the experimental findings are in contradiction with this hypothesis as the characteristic frequency, f i, is higher and the transfer resistance, Rti, is lower for Li+ compared with the values corresponding to Li+·H2O. This observation is likely due to the selectivity of the Li-birnessite type MnO2 toward Li+ ions. The kinetics parameters and the transfer resistance values of both Li+ species vary slightly with the applied potential (Figure 4a,c). This dependence implies their faradaic contribution to charge storage in LiClO4 medium. The capacitance values related to each Li+ species calculated from ac-electrogravimetry data corroborate this slight dependence on the applied potential (Table S2 in Supporting Information). Nevertheless, because of its faster kinetics, Li+ can be associated with the electroadsorption process (reaction mechanism I in the Introduction) and Li+·nH2O to the redox process (reaction mechanism II in the Introduction). Water-transfer behavior in LiClO4 is different compared to the measurements in NaClO4 media. The water flux direction changes with the potential applied (Table S1 in Supporting

CONCLUSIONS The results presented in this study on Li-birnessite type MnO2 thin films establish that the pseudocapacitive charge-storage mechanisms can be monitored by ac-electrogravimetry. The technique provides the exact identification and kinetic information on the charged and uncharged species, which is not possible with classical EQCM measurements. Pseudocapacitive behavior of Li-birnessite type MnO2 thin films in LiClO4 and NaClO4 aqueous solutions involves Li+ ions, Na+ ions, and their respective hydrated ionic species, as detected by our acelectrogravimetry study. This indicates that there is a population of hydrated Li+ or hydrated Na+ ions losing their hydration shell before being transferred at the electrode/ electrolyte interfaces. The number of water molecules in the hydration shell of the hydrated ions, the kinetics (fci), and resistance (Rti) values of charged and noncharged species transferred at the electrode/electrolyte interfaces were estimated from our ac-electrogravimetry study. The opposite flux direction of free water molecules was also detected by acelectrogravimetry. Either water molecules (in the interlayer distances of the material) liberate space for cations or water molecules due to the partial desolvation of hydrated cations are expelled from the film so that partially hydrated cations can access to less attainable sites. To our knowledge, this is the first time that a fair and in situ determination of this effect of desolvation has been clearly demonstrated. Preliminary results in nonaqueous media (LiClO4 or NaClO4 in acetonitrile electrolyte) also indicate the presence of two different cationic species, Li+ ions and Na+ ions and their respective solvated species, as well as free solvent molecules. This implies that, similar to the results in aqueous media, there is a population of solvated Li+ or solvated Na+ ions losing their hydration shell before being transferred at the electrode/electrolyte interfaces. These results are exciting, and the combination of fast QCM with electrochemical impedance spectroscopy offers a great opportunity for a better understanding of the ion transfer and extent of solvation/desolvation in porous and/or layered materials which is not possible with single EQCM measurements. Therefore, the establishment of the ac-electrogravimetry characterization in the energy materials domain is significant for designing optimized materials for the next-generation high energy density pseudo- and/or supercapacitors.

■

26558



■

Article

(14) Gabrielli, C.; Garcia-Jareno, J. J.; Keddam, M.; Perrot, H.; Vicente, F. Ac-electrogravimetry study of electroactive thin films. I. Application to prussian blue. J. Phys. Chem. B 2002, 106, 3182−3191. (15) Sel, O.; To Thi Kim, L.; Debiemme−Chouvy, C.; Gabrielli, C.; Laberty-Robert, C.; Perrot, H.; Sanchez, C. Proton insertion properties in a hybrid membrane/conducting polymer bilayer investigated by AC electrogravimetry. J. Electrochem. Soc. 2010, 157, F69−F76. (16) To Thi Kim, L.; Sel, O.; Debiemme−Chouvy, C.; Gabrielli, C.; Laberty-Robert, C.; Perrot, H.; Sanchez, C. Proton transport properties in hybrid membranes investigated by ac-electrogravimetry. Electrochem. Commun. 2010, 12, 1136−1139. (17) Nakayama, M.; Kanaya, T.; Lee, J. W.; Popov, B. N. Electrochemical synthesis of birnessite-type layered manganese oxides for rechargeable lithium batteries. J. Power Sources 2008, 179, 361− 366. (18) Pinaud, B. A.; Chen, Z.; Abram, D. N.; Jaramillo, T. F. Thin films of sodium birnessite-type MnO2: Optical properties, electronic band structure, and solar photoelectrochemistry. J. Phys. Chem. C 2011, 115, 11830−11838. (19) Sauerbrey, G. The use of a quartz crystal oscillator for weighing thin layers and microweighing applications. Z. Phys. 1959, 155, 206− 222. (20) Mahlar, J.; Persson, I. A. Study of the hydration of the alkali metal ions in aqueous solution. Inorg. Chem. 2012, 51, 425−438. (21) Ravinder, N. R.; Ramana, G. R. Sol−gel MnO2 as an electrode material for electrochemical capacitors. J. Power Sources 2003, 124, 330−337.

ASSOCIATED CONTENT

S Supporting Information *

Electropotentiostatic deposition of Li-birnessite type MnO2 coatings and their structural characterization via FEG-SEM, EDX, XRD, and XPS analyses; the experimental setup of acelectrogravimetry and the details of the theoretical model considering the transfer of two cations and free solvent molecules at the electrode/electrolyte interfaces. This material is available free of charge via the Internet at http://pubs.acs.org.

■

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS “Le labex MATISSE-Sorbonne Universités UPMC Univ Paris 06” is highly appreciated for the financial support of “Master 2” internship of C. R. A. Ms. Françoise Pillier and Mr. Cyrille Bazin are acknowledged for their help with the FEG-SEM and XRD measurements, respectively.

■

REFERENCES

(1) Conway, B. E. Electrochemical Supercapacitors: Scientific Fundamentals and Technological Applications; Kluwer: New York, 1999. (2) Simon, P.; Gogotsi, Y. Materials for electrochemical capacitors. Nat. Mater. 2008, 7, 845−854. (3) Tarascon, J. M.; Armand, M. Issues and challenges facing rechargeable lithium batteries. Nature (London, U.K.) 2001, 414, 359− 367. (4) Whittingham, M. S. Lithium batteries and cathode materials. Chem. Rev. (Washington, DC, U.S.) 2004, 104, 4271−4301. (5) Toupin, M.; Brousse, T.; Belanger, D. Charge storage mechanism of MnO2 electrode used in aqueous electrochemical capacitor. Chem. Mater. 2004, 16, 3184−3190. (6) Brezesinski, K.; Wang, J.; Haetge, J.; Reitz, C.; Steinmüller, S. O.; Tolbert, S. H.; Smarsly, B. M.; Dunn, B.; Brezesinski, T. Pseudocapacitive contributions to charge storage in highly ordered mesoporous group V transition metal oxides with iso-oriented layered nanocrystalline domains. J. Am. Chem. Soc. 2010, 132, 6982−6990. (7) Kanoh, H.; Tang, W.; Makita, Y.; Ooi, K. Electrochemical intercalation of alkali-metal ions into birnessite-type manganese oxide in aqueous solution. Langmuir 1997, 13, 6845−6849. (8) Kuo, S.-L.; Wu, N.-L. Investigation of pseudocapacitive chargestorage reaction of MnO2·nH2O supercapacitors in aqueous electrolytes. J. Electrochem. Soc. 2006, 153, A1317−A1324. (9) Tsai, W.-Y.; Taberna, P. L.; Patrice, S. Electrochemical quartz crystal microbalance (EQCM) study of ion dynamics in nanoporous carbons. J. Am. Chem. Soc. 2014, 136, 8722−8728. (10) Chmiola, J.; Yushin, G.; Gogotsi, Y.; Portet, C.; Simon, P.; Taberna, P. L. Anomalous increase in carbon capacitance at pore sizes less than 1 nm. Science 2006, 313, 1760−1763. (11) Levi, M. D.; Salitra, G.; Levy, N.; Aurbach, D.; Maier, J. Application of a quartz-crystal microbalance to measure ionic fluxes in microporous carbons for energy storage. Nat. Mater. 2009, 9, 872− 875. (12) Barisci, J. N.; Wallace, G. G.; Baughman, R. H. Electrochemical quartz crystal microbalance studies of single-wall carbon nanotubes in aqueous and non-aqueous solutions. Electrochim. Acta 2000, 46, 509− 517. (13) Chu, Y.-H.; Hu, C.-C.; Chang, K.-H. Electrochemical quartz crystal microbalance study of amorphous MnO2 prepared by anodic deposition. Electrochim. Acta 2012, 61, 124−131. 26559


New Insights into Pseudocapacitive Charge-Storage Mechanisms in Li

Recommend Documents