J. Phys. Chem. B 1999, 103, 1499-1508
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Dilute Graphite-Sulfates Intercalation Stages Studied by Simultaneous Application of Cyclic Voltammetry, Probe-Beam Deflection, In situ Resistometry, and X-ray Diffraction Techniques M. D. Levi, E. Levi, Y. Gofer, and D. Aurbach* Department of Chemistry, Bar-Ilan UniVersity, Ramat-Gan 52900, Israel
E. Vieil and J. Serose Laboratoire d’Electrochimie Moleculaire, UMR 5819(CEA + CNRS + UniVersite J. Fourier), Departement de Recherche Fondamentale sur la Matiere Condensee, CEAsGrenoble, 38054 Grenoble Cedex 9, France ReceiVed: July 31, 1998; In Final Form: NoVember 23, 1998
This paper describes the stoichiometry, kinetics, and electronic resistivity of dilute graphite-sulfates intercalation compounds (stages pristine toVIII) obtained during electrochemical oxidation of graphite in 4M H2SO4 solution. The stoichiometry of the reaction was determined by the use of cyclic voltammetry and in situ XRD. Dilute graphite-sulfates phases are formed continuously as solid solution in the graphite matrix. The electronic resistivity of the composite graphite electrode as a function of potential depends largely on parasitic reactions, i.e., partial oxidation of graphite particles. The ionic flux of the HSO4- species across the electrode-solution interface was traced by deflection of a probe-laser beam (“mirage”, or PBD technique) measured simultaneously with cyclic voltammetry. Convolution of the resulting deflectograms, measured at different distances from the electrode, allowed the calculation of diffusion coefficient of the active species in solution. Basic electroanalytical features of Li-ion intercalation into graphite from Li+-containing aprotic electrolytes are compared with those for HSO4- in aqueous 4M H2SO4.
Introduction Well-ordered graphite is known to be one of the important layered materials which readily reacts with various electron donors (i.e., alkali and alkaline earth metals, lanthanides, etc.) and electron acceptors (halogens, acidic oxides such as N2O5 and SO3, and Bro¨nsted acids: HF, HNO3, H2SO4, and oleum).1-6 The above species penetrate into the graphite host matrix during a chemical or an electrochemical reaction. Depending on the intercalation level, the accommodation of guest species at their host matrix sites occurs, usually, via formation of a series of definite staged phases. This is evident from the change in the position of the 00l XRD peaks with the intercalation level (these peaks are the strongest in the powder XRD patterns of graphite).2 The mechanisms of the electrochemical intercalation are usually more complicated as compared to those of the chemical intercalations since the former relates to several components of the system (solvent, salt, current collector, etc.) which can interfere with the major process. Generally, electrochemical intercalation consists of two coupled processes: insertion of an ionic intercalant and a change in the host’s electronic structure. The flux of ionic species during intercalation can be monitored using electroanalytical techniques specially designed for this, for example, probe beam deflection (or “mirage”) technique (PBD)7-12 and electrochemical quartz-crystal microbalance (EQCM).12-15 Moreover, application of different techniques, such as in situ resistometry, may result in better characterization of the flux of electronic species. Thus, these two groups of in situ techniques are complementary with regard to their application for studying intercalation mechanisms. This * To whom correspondence should be addressed.
was nicely demonstrated in the case of thin films of electronically conducting polymers. It is of interest to extend the application of the above in situ techniques for the characterization of insertion reactions into composite powder electrodes that are widely used both as anodes and cathodes in high-energy-density rechargeable Li-ion and other types of batteries. Improvement of their performance cannot be reached without a proper understanding of the intercalation and capacity-fading mechanisms as well as the nature of the composition-structure relationship. Combined application of cyclic voltammetry (CV), PBD, in situ resistometry, and XRD may be very useful in this respect. Of course, utilization of these techniques is the most effective if powder electrodes are prepared as relatively thin coatings on an inert support. We have reported on the preparation of such electrodes, of several micrometers thick, which are composed of 90% of synthetic graphite powder and 10% of polyvinylenedifluoride (PVdF) binder.16-18 These thin-coated electrodes were used for detailed electroanalytical characterization of the Li-ion insertion process from aprotic solutions. The very low intercalation potentials typical for the above reaction (ranged from 0.3 to 0 V vs Li/Li+) make it a prerequisite that all sample manipulations will be performed in an inert atmosphere glovebox. Since it was very difficult to perform PBD experiments in a glovebox with the current apparatus, we gave preference to another insertion process as a model, in which the composite graphite electrode was intercalated with HSO4- in aqueous H2SO4 solutions. In the past years, this reaction has attracted much attention, with the main interest focused on staging phenomena and intercalation kinetics of the graphite in contact with concentrated H2SO4. Cyclic voltammetry,5 electrochemical
10.1021/jp9832443 CCC: $18.00 © 1999 American Chemical Society Published on Web 02/17/1999
1500 J. Phys. Chem. B, Vol. 103, No. 9, 1999 impedance spectroscopy,20 chronopotentiometry coupled with dilatometry (the latter technique probes the changes in the electrode’s thickness as a function of the intercalation level), and in situ XRD2 were the primary tools for these studies which were performed mainly with HOPG electrodes. The purpose of this work was to characterize the intercalation mechanism and side anodic reactions occurring in practical composite graphite electrodes during their cycling in dilute H2SO4 solution utilizing simultaneously CV, PBD, in situ resistometry, and XRD. Experimental Section Thin-coated graphite electrodes for CV measurements were prepared as described previously.16-18 They were several micrometers thick covering a 1.2 × 1.2 cm2 Pt plate. A matched Pt or Pb plate was used as a counter electrode, whereas silver wire positioned close to the working electrode in the same solution served as a quasi-reversible reference electrode. H2SO4 (4 M) used in this work was obtained by the dilution of 96% analytical H2SO4 with an appropriate amount of doubledistilled water. Cyclic voltammetry and electrochemical impedance measurements were performed at ambient atmospheric conditions. The corresponding instrumentation was described in details elsewhere.16-18 PBD experiments were performed as already described for the doping of thin films of electronically conducting polymers.10 The diameter of the laser beam spot in the focus was close to 80 µm. Although this diameter has a finite size, the measurement of the concentration gradient can be considered as a localized determination. The reason for this is due to the fact that the beam intensity is not uniform on a cross section and presents a bell-shaped profile (Gaussian distribution). It has been theoretically shown19 that this finite-size effect is negligible. Experimentally, it has already been shown that the PBD technique could clearly identify exclusively monoionic exchange with an electroactive film and measure diffusion coefficients in solution.10 The deviation of the beam under cycling of potential was measured at several distances between the electrode surface and beam’s center, starting from a distance of the closest approach (80 µm). This distance is to be compared with the thickness of the studied graphite coatings, estimated to be several micrometers thick (the expected addition of 10-15% thickness due to swelling is negligible). Mathematical treatment of the PBD data including the convolution procedure has been previously reported.10 The current collector in these measurements was a Pt disk (4 mm in diameter) hermetically pressed into a glass tube. A drop of slurry containing KS-6 powder (0.4-0.5 mg) and PVdF binder in 1-methyl-2-pyrrolidone was placed on the Pt surface and then dried in an oven at 120 °C for at least 20 min. In situ XRD was performed in a three-electrode cell using the instrumentation, which has already been described.18 In this case, we used a mixture of graphite powder and PVdF (about 20 mg of active mass of the same composition) pressed onto a Pt net (1.6 cm2). The cell was configured with a circular window in the front of a polyethylene cylindrical cell and hermetically sealed with a thin polyethylene film as a window. The electrode’s surface was kept as close to this window as possible in order to minimize the thickness of the solution layer between them. This layer considerably decreases the intensity of the XRD response and hence reduces the resolution of the data obtained. In situ resistometry has recently been introduced in our laboratory at CEAsGrenoble in order to measure the electronic conductivity of thin films of conjugated polymers as a function
Levi et al. of the applied potential. The setup consists of an array of six parallel microelectrodes on which the electroactive film is deposited. The two outmost electrodes (2 and 6) connected together to the potentiostat for the voltammetry (working electrode). For the resistommetry measurements, a dc current bias is applied between the two middle electrodes (2 and 4). The current is applied by means of a floating source (a 10.5 V battery and a potentiometric resistor for varying the voltage and the current). The sample preparation consists of an application of a small drop of graphitesPVdF slurry in 1-methyl-2pyrrolidone that was placed onto this microelectronic “comb”. The setup designed as an array of thin gold strips of 100 µm width with a 20 µm gap between them. This coating was polarized in a three-electrode cell with a platinum plate counter electrode. Silver wire served as a reference electrode in all cases. The absolute values of resistance of the as-prepared graphite coatings polarized in the sulfuric acid solution were obtained by the calibration of the apparatus with a series of standard resistances. Results and Discussion 1. Cyclic Voltammetry and Electrochemical Impedance Characterization of the Graphite Electrode during Oxidation in Dilute H2SO4. Figure 1a shows the first five cyclic voltammograms of a thin graphite electrode in 4 M H2SO4. The scans were measured at a rate of ν ) 10 mV s-1 in the potential range from the OCV of the fresh electrode, ∼0.3 V (vs Ag wire) up to 1.4 V (the range 0.3-0.8 V is not shown in the figure since within this potential range the measured current was negligible). A progressive increase of the reversible capacity of the graphite electrode on cycling is clearly seen from this figure. After 10 cycles, the CV response has reached a stable reproducible value. This allowed to further cycle the electrode at different scan rates ranged from 1 to 50 mV s-1. The corresponding voltammograms are shown in Figure 1b (the currents were normalized with respect to the scan rate applied). It is seen from this figure that the voltammograms possess a small but clearly resolved anodic peak at 1.23 V and a broad cathodic peak centered at 1.16 V on the reverse scans with a poorly resolved shoulder at 0.92 V. The separation between the anodic and cathodic peaks becomes somewhat larger with the increase of scan rate as expected for a kinetically limited intercalation reaction. Diffusion control of the intercalants in the graphite particles is hardly attained in the range of ν used since the current response is proportional to ν rather than to ν1/2. A typical charge vs potential curve calculated from the cyclic voltammogram measured at ν ) 10 mV/s is shown in Figure 2. A sloping shoulder at 1.28 V reflects the smaller intercalation peak in the CV at 1.23 V. It is clear from the comparison of Figures 1a and 2 that the second (large) anodic peak is either superimposed on that of the lower intercalation stages (more advanced) or some side reactions that prevent the electrode charging from eventual saturation. As seen from Figure 2, the intrinsic hysteresis between the charge and discharge is close to 0.2 V in average. The existence of superimposed side reactions on the intercalation process can be ascertained as follows: (i) by observation of the decrease in the effective Coulombic efficiency of the overall reactions as the scan rate becomes smaller, (ii) by analyzing the change in the shape of the complex-plane impedance spectra (Nyquist plots) with increasing the electrode potential, (iii) by referring the electrode’s phase composition
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Figure 2. A plot of charge and discharge capacity vs potential obtained by integration of the cyclic voltammetry curve (ν ) 25 mV s-1).
Figure 1. (a) Cyclic voltammograms measured subsequently with a thin-coated powder graphite electrode (∼1 mg active mass with 10% PVDF binder on a 1.2 × 1.2 cm2 Pt foil) in 4 M H2SO4. Scan rate ν ) 10 mV s-1. Cycles numbers are indicated. (b) Cyclic voltammograms (normalized with respect to the scan rate) measured subsequently with thin-coated powder graphite electrode in 4 M H2SO4 solution at ν ranged from 1 to 50 mV s-1, as indicated.
derived from the in situ XRD to that obtained from the coulometric curves. When comparing the ratio between the areas under the cathodic and anodic branches of the voltammetric curves shown in Figure 1b, a decrease in the effective Coulombic efficiency becomes obvious with a decrease in the scan rate. Thus point i supports the involvement of some side reactions to the process. Figures 3a and b present Nyquist plots of an electrode in the potentials range from 1.21 to 1.4 V. Figure 3a covers the frequencies range 25 kHz to 5 mHz, whereas Figure 3b shows the high-frequency domain only. As seen from the latter figure this domain consists mainly of a slightly depressed semicircles (19.5° depression with respect to the Z′-axis) of surprisingly small diameter, 0.67 Ω (1 mg active graphite mass). This is by 1-2 orders of magnitude smaller than that observed for Li-ion intercalation into graphite (of similar electrode’s mass). Nyquist plots shown in Figures 3a and b can be interpreted within the framework of Voigt-Frumkin and Melik-Gaykazyan imped-
Figure 3. A family of Nyquist plots measured with thin graphite electrode as a function of potential for the ac frequency ranged from (a) 25 kHz to 5 mHz and from (b) 25 kHz to 63 Hz.
ance (a finite-space diffusion model) which has been shown to be very useful for the analysis of impedance spectra of lithiated graphite16 and lithiated transition metal oxides.20 We thus treated the high-frequency semicircle as a superposition of two semi-
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circles: one reflects the presence of some surface species covering the graphite particles whereas the other is linked to the HSO4- interfacial charge-transfer reaction. The right-hand side of the Nyquist plots of Figure 3b (i.e., as the frequency decreases) is characterized by a very narrow region identified with Warburg impedance. Note that both the two domains, the semicircle and the Warburg one, appear also on the Nyquist plots of LixC6 and LixCoO2 intercalation electrodes.20 However, in contrast to these electrodes, in the low-frequency domain (Figure 3a) the tendency of reaching a pure capacitive behavior (at the beginning of intercalation) is distorted and an evolution of a new semicircle becomes clearly resolved at higher electrode potentials. Tentatively, we ascribe this evolving second semicircle to side reactions. Our assumption is based on the following reasoning: (i) impedance spectra of graphite electrode in contact with concentrated H2SO4 solutions revealed a limiting capacitive domain at low frequencies.21 It is known that the effect of side reactions on the oxidation of graphite in highly concentrated solutions of H2SO4 is much smaller than that in diluted solutions.4,22 (ii) Nyquist plots of LixC6 electrode in aprotic solutions showed a limiting capacitive behavior with only small deviation from a semicircle of a very high diameter.16 Similar behavior was observed also for lithiated transition metal oxides.20 The common electroanalytical feature of these two systems related to the problem under study is the contribution of possible side reactions to the overall charge and discharge processes. In these cases the contribution of parasitic reactions is negligible because of effective surface passivation in the former case, and the high anodic stability of the solvents and electrolytes used, in the latter one. It is known that modeling an impedance response may be misleading since different electric circuits may fit any Nyquist plot. In this case, the comparison between the calculated parameters obtained from the modeling and that obtained by other techniques, may be crucial in identifying an adequate model. This will be further illustrated by calculation of the chemical diffusion coefficient, Dc, of HSO4- anions in the graphite matrix. Assuming that the initial part of the sloping line of the Nyquist plots (immediately after crossing the highfrequency semicircle toward the lower frequencies) is, in fact, the true Warburg region, one can calculate Dc according to the following formulas:17,23
τ ) l2/Dc ) 2 (Qint Aw dX/dE)2
(1)
Where τ is the characteristic diffusion time of HSO4- ions in the graphite platelet particles and l is the characteristic diffusion length roughly estimated as half of the average particle size; the Warburg slope Aw ) ∆Re/∆ω-1/2 ) ∆Im/ ∆ω-1/2 (∆Re and ∆Im are the differences between the real and the imaginary components of the impedance, corresponding to a finite variation in the frequency), Qint is the maximum intercalation charge, X is the intercalation level, and thus Qint dX/dE represents the differential capacity. Using the above formulas, it was found that, at the beginning of the insertion process (E ) 1.201 V), Dc is 1.59 × 10-5 cm2 s-1, whereas at higher level of insertion (E ) 1.305 V) Dc ) 2.0 × 10-6 cm2 s-1. The first of the two values is close to the literature data, 1.0 × 10-5 cm2 s-1 24, although values as high as 1.7 × 10-4 cm2 s-1 were reported for more concentrated graphite sulfates.20,25 It is possible that the difference in Dc obtained by different authors originates from the uncertainty in values of l, thus the comparison between τ could be more reasonable. Note also that, as the intercalation level increases,
Dc tends to decrease, i.e., this behavior is similar to that for diluted phase 1 of lithiated graphite.17,18 However, the absolute values of Dc for the latter compounds are by 4-5 orders of magnitude lower than that for the graphite sulfate under study. 2. Simultaneous CV and in Situ XRD Studies of the Graphite Sulfate Intercalation Compound. For the determination of the graphite sulfate stoichiometry, CV and in situ XRD were measured simultaneously with thicker electrodes containing 16 mg of the active graphite mass. It was found that similar to the case of the first cycles with thin electrodes, the fraction of the graphite’s active mass utilization increased with cycle’s number but at a slower rate. To reach the steady-state chargedischarge regime rapidly, the electrode was cycled at ν ) 5 mV s-1 from 0.4 up to 1.4 V. Prior to the reverse sweep, the potential was held at the upper potential limit for different periods of time ranged from 0 to 90 min. Figure 4a shows a dramatic increase in the reversible intercalation capacity with the delay time which probably can be connected with a slow rate of wetting the porous electrode’s active mass containing hydrophobic graphite and PVdF binder. Note that the voltammetric curves measured with thick electrodes show the same two peaks which have appeared in the CV of a thinner electrode, a behavior that supports the assumption that the thick electrodes are behaving in a manner similar to the thin one. To study the nature of the reactions occurring in the anodic potentials limit, the upper vertex potential was shifted stepwise from 1.4 to 1.6 V holding at the limit for 10 min prior to the reverse scan. Figure 4b summarizes the results obtained. The second (the most positive) voltammetric peak becomes very broad and hence the charge and discharge capacities related to this peak drastically increase with the anodic potential limit. Polarization at 1.6 V seems to be critical to the electrode’s behavior since the reverse sweep (curve 4 in Figure 4b) shows a pronounced shift in peak potentials which results in the increase of the peak potential separation. Note that the subsequent anodic potential sweeps caused a further increase in current as compared to that for the preceding sweeps (compare Figures 4b and c). Curve 1 in Figure 4c was measured at ν ) 5 mV/s in the potential range 0.4 to 1.55 V holding at the vertex potential for 10 min. The subsequent potential scan was performed down to -0.3 V, followed by constant polarization at this potential for 1h. The integration of current with respect to time at each separate charge and discharge steps allowed a semiquantitative calculation of the extent of various side anodic reactions. One possible side reaction can be oxygen evolution on the electrode. Figure 4c shows a voltammetric curve of a bare Pt electrode with characteristic irreversible oxygen evolution reaction. Oxidation of HSO4- anions to S2O82- at these conditions can hardly take place since the cathodic branch of curve 2 in Figure 4c shows no appreciable current down to -0.3 V, at which point the generated persulfate electroreduction is expected to occur. In contrast, the cathodic branch of the voltammetric curve of the graphite electrode consists of the broad peak around 1.16 V which has been previously assigned to the HSO4- deinsertion reaction and another pronounced peak around of -0.1 V. It was suggested that this peak could be due to the reduction of S2O82- anions obtained by oxidation of HSO4- anions within the graphite matrix. Since such an oxidation reaction does not occur at a bare Pt electrode, it is reasonable to assume that the graphite host matrix provides a more facile oxidation medium for the inserted HSO4- species. This result can be considered as a manifestation of the electrocatalytic capability of graphite, which agrees well with
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Figure 4. (a) Cyclic voltammograms measured subsequently with a thick powder graphite electrode (∼16 mg active mass) in 4 M H2SO4. Scan rate ν ) 5 mV s-1. Cycle numbers are indicated. Prior to the reverse cathodic scan, the electrode was polarized at the upper limit potential for different time (min): (curve 1) 20, (curve 2) 40, (curve 3) 60, (curve 4) 90. (b) Cyclic voltammograms measured subsequently with thick powder graphite electrode (∼16 mg active mass) in 4 M H2SO4. Scan rate ν ) 5 mV s-1. The upper limit potential was as follows: (curve 1) 1.38, (curve 2) 1.46 V, (curve 3) 1.50 V, (curve 4) 1.60 V. Prior to the reverse cathodic scan, the electrode was polarized at the upper vertex potential for 10 min. (c) Cyclic voltammograms measured with thick powder graphite electrode (1) and Pt foil (2) in 4M H2SO4. Curve 3 shows a cyclic voltammogram obtained with Pt in 4M H2SO4 with 0.002 M K2S2O8. All three curves were obtained at ν ) 5mV s-1. Prior to the reverse cathodic scan, the electrodes were polarized at the upper vertex potential for 10 min. At the end of the cathodic scan, the electrodes were potentiostatically polarized at -0.3 V for 1 h.
the similar early suggestion made by Metrot et al. concerning the oxidation of graphite in concentrated sulfuric acid.6 Persulfate formation was further proved by a direct observation of a peak assigned to S2O82- reduction on a bare Pt electrode upon addition of K2S2O8 to the dilute sulfuric acid solution. It was found that the height of this peak was proportional to the concentration of S2O82- in the solution. The low potential of S2O82- reduction peak is not surprising since this reaction is known to proceed at a very high overvoltage.26 The values of the current efficiencies for the intercalationdeintercalation processes were calculated as follows: during the anodic sweep, 1.5 C of charge passed. In addition, 2.5 C passed during the potentiostatic polarization of the electrode at 1.55 V. Thus the total anodic integrated charge was 4.0 C. From the integration of the corresponding peaks, the deinsertion of bisulfate anions corresponds to -0.73 C, whereas persulfate
reduction involved -2.57 C. The rest amount of charge, i.e., -(4.0-3.3) ) -0.7 C, can be ascribed mainly to irreversible oxygen evolution at high anodic potentials and, to a lesser extent, to the overoxidation of the graphite. Slow reduction of the graphite’s surface oxidation products can be traced as an increase in the ”background” current between the potentials corresponding to the HSO4- deinsertion and persulfate reduction reactions (see Figure 4c). The total discharge capacity 0.7 C from deinsertion reaction results in a specific value of 12.7 mAh g-1. Thus, the stoichiometry of the insertion/deinsertion reaction to/from the graphite in diluted sulfuric acid solutions is expressed by the formulas C160 HSO4, which is very close to the VIII stage, C192HSO4. As opposed to the intercalation reaction from a concentrated H2SO4, higher concentrated stages cannot be obtained in the diluted acid.
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Figure 5. Weight fraction of the dilute graphite sulfate phase vs potential derived from in situ XRD. The scan rate used as indicated.
Figure 6. In situ XRD patterns (00l peak) of the dilute graphite sulfate phase as a function of potential as indicated. The measurements were taken after 0.5 h potentiostatic polarization at the above potentials.
This conclusion was verified using in situ XRD. Figure 5 shows the decrease in the intensity (arbitrary units) of the 002 peak of the pristine graphite during potential sweeping at three different rates. These curves are surprisingly similar to the corresponding charge-discharge ones (compare Figures 5 and 2). Interestingly, two separate anodic peaks on the cyclic voltammogram of graphite result in small but distinct changes in the slope of the phase composition-potential curves. Figure 5 shows unambiguously that the amount of the graphite in the pristine phase decreases with potential during intercalation. Figure 6 demonstrates that the formation of the CxHSO4 intercalation compound is reflected by a gradual shift of the 00l peak position toward lower angle values. Polarization of the electrode at 1.32-1.33 V results in 00l peak shift by 0.96° which corresponds to the 00l peak position dn ) 3.47 Å(calculated from Bragg’s equation):
dn ) [di + 3.35(n - 1)]/l
(2)
where di is the distance between two graphitic planes on either
Figure 7. (a) Simultaneously measured cyclic voltammogram (ν ) 20 mV s-1) and the probe beam deflectogram obtained at a x ) 230 µm distance from the electrode surface. (b) Comparison between the experimental deflectogram shown in a and the current projected at the distance x ) 230 µm by convolution.
side of the intercalation layer (equal to 7.98 Å for the bisulfate anion1), n is the stage number, and l ) (n - 1) + 2. In the case where n ) 8 in eq 2, dn value obtained is 3.49 Å, thus the intercalation compound obtained at 1.32 V is almost pure phase VIII. This conclusion is in line with Figures 1a and b: the potentiostatic polarization at this potential allows for attaining the discharge capacities of 9-10 mAh/g which is close to 11.6 mAh/g expected for the VIII stage. 3. PBD Characterization of a Thin Graphite Electrode during Intercalation with HSO4- Anions. PBD allows to observe directly variations in the concentration gradient of the electroactive species in the solution phase in proximity to the surface of the electrode. Figure 7a shows simultaneously measured cyclic voltammogram (ν ) 20 mV/s) and the corresponding beam deflection (θn mrad-1. n is the refractive index of the solution) at a distance of x ) 230 µm from the electrode surface. The measured current i is proportional to the flux of the species under consideration, Ji, close to the
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Figure 8. Comparison between the experimental deflectogram measured at the distance of the closest approach to the electrode surface, x ) 80 µm and the deviation at x ) 230 µm projected at the distance x ) 80 µm by convolution.
electrode-solution
interface:10
∑ ziJi(0,t)
i(t) ) FA
(3)
where A is the surface area of the electrode and z is the ionic charge. The deviation of the beam θ, initially aligned in parallel to the electrode’s surface at a certain distance x from it, is proportional to the sum of the fluxes of solution species (anions and cations):
θ(x,t) ) AL (-hC JC (x,t) - hA JA (x,t))
(4)
with hi ) (1/n)(dn/dc)(1/2Di) where L is characteristic length reflecting the distance along which the refraction index is perturbed, c is the concentration of species in solution near the electrode interface, and Di is the diffusion coefficient of the species as measured at the interface with the solution. The relation between the deviation θ (eq 4) and the current (i) (or the flux J(t)), in eq 3 is given by the convolution product:
θ(x,t) ) -(h/zFA)F(x,t)J(t)
(5)
where the convolution transfer function has a form:
F(x,t) ) [x/2(πDt3)1/2] exp (-x2/4Dt)
(6)
Here D is the mean diffusion coefficient of both cation and anion, since the electroneutrality in solution requires equality of concentrations of oppositely charged species. Using a homemade software elaborated at the CEA-Grenoble the current can be easily convoluted into θ(x,t) using eqs 3-6. Figure 7b compares the experimental nθ vs E plot, measured at a distance x ) 230 µm, with that convoluted from the cyclic voltammogram shown in Figure 7a. It is clearly seen that the two curves are in good agreement. Figure 8 compares the deflection measured at x ) 80 µm with the one at x ) 230 µm, projected, by convolution, to the distance x ) 80 µm. Once again, a similarity between the two curves is obviously seen. Figure 9 shows a plot of the convolution parameter (derived from the fitting procedure) vs the relative distance of the beam
Figure 9. Plot of the convolution parameter, xD-1/2 vs the distance between the electrode surface and the probe beam spot.
from the electrode surface. From this figure, the effective ionic diffusion coefficient of the intercalants in the solution was calculated to be D ) 2.3 × 10-3 cm2 s-1. This value is about 2 orders of magnitude higher than the one typically obtained for various species in aqueous solutions. The specific conductivity of concentrated H2SO4 is very high, 0.1 S cm-1 25 from which the effective diffusion coefficient, of the order of 10-4 cm2 s-1, can be derived. The reason for the abnormal high conductivity is that both H+ and HSO4- are able to diffuse via an exchange mechanism. In our case, we may also expect rather high value of D as compared to other electrolytic solutions. The rest of the discrepancy between the measured and expected D may be indicative of a more complicated intercalation mechanism. A higher diffusion coefficient than expected may result from a higher actual flux than that of anions alone The increase of the flux may originate from the cointercalation of neutral molecules of H2SO4 together with HSO4- anions. In a given concentration gradient, higher flux implies a higher effective diffusion coefficient. The ratio ([HSO4-]/[H2SO4]) between the intercalating species in concentrated sulfuric acid solution was found to be 2.5-3.0.2,3 However, according to Metrot et al.,22 in 4 M H2SO4, cointercalation of H2O with a ratio value [H2O]/[HSO4-] ) 9.6 is expected rather than the cointercalation of H2SO4. Thus, in general case, the insertion reaction during electrochemical oxidation of graphite in H2SO4 may be expressed by the following equation:22
xC + HSO4- + yH2SO4 + zH2O ) Cx + HSO4- yH2SO4zH2O + e- (7) where y and z depend mainly on the concentration of H2SO4 in solution. The case of intercalation reaction for which y and z take nonzero values needs a special consideration for the quantitative treatment of the PBD response (this is beyond the scope of the present work). 4. In Situ Resistometry during Intercalation of Graphite with H2SO4. In situ electronic resistivity of the graphite composite electrode was measured at three different scan rates: 5, 10, and 20 mV/s simultaneously with the cyclic voltammetry, see Figures 10a, b, and c, respectively. As seen in Figure 10a,
1506 J. Phys. Chem. B, Vol. 103, No. 9, 1999 Figure 10. Cyclic voltammograms and in situ resistometric curves measured simultaneously at the scan rates: (curve a) 5, (curve b) 10, and (curve c) 20 mV s-1. Latin letters from a to g designate separate specific regions of these curves.
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Dilute Graphite-Sulfates Intercalation Stages resistometric curves reveal several characteristic points during oxidation (a, b, c) and reduction (d, e, f, g). The initial potential sweep from 0 V to point a results in a moderate and monotonic decrease in the resistance of the composite electrode. The extent of the insertion at this stage is very small as compared to that of the total insertion process. The beginning of the intercalation reaction (point b) causes small but distinct increase in the resistance. Further intercalation up to point c gives a drastic increase of resistance which continues up to point d (at that point the current changes its sign). Surprisingly, the resistance continues to increase up to point e at which point the current passes through its major deintercalation peak. In contrast to the expected pronounced changes in the host during the deintercalation, in the potential region between points e and f, the resistance remains practically unchanged. At this latter point f (0.52 V) the resistance drastically drops down and then decreases moderately upon approaching the final potential 0 V. Similar characteristic points can also be observed on the resistometric curves measured at higher scan rates (Figure 10 b and c). It is seen that the higher the scan rate the lower is the maximum variation in resistance. Comparison between Figures 10, 2, and 5 shows that the hysteresis in the resistometric curves is much higher than that on the charge-discharge curves or the curves characterizing the variation of the amount of the pristine graphite phase determined by XRD. This clearly demonstrates that the intercalation-deintercalation reaction can only partially influence the electronic conductivity of the graphite electrode. One has to bear in mind that the electrode under study is a composite powder electrode possessing partially oriented platelet particles. Taking this into consideration, specific regions on the resistometric curves can be assigned as follows. Upon oxidation (a), A decrease in resistance is mainly connected with the increase of in-plane conductivity of graphite particles. For HOPG crystals, the slope of the plot of conductivity vs intercalation level has been shown to take rather high positive values.27 (b and c) A drastic increase of the resistance is due to superposition of various processes. Advancement of the intercalation of bisulfate anions into graphite should result in the decrease of in-plane resistivity of the individual particles (as is the case of HOPG), but this effect is completely screened by the oxidation of carbon atoms on the particles surface and by adsorption of oxygen on the surface of the current collector. When the last two factors are dominating it may lead to considerable increase in the interparticle and interface resistivities. Upon reduction (d) the increase in resistivity is still controlled by the two latter factors (the current remains anodic in the c-d sweep region). (e-f) the behavior is reversed with respect to that at point a: the intensive deintercalation causes increase in resistivity. (f-g) Oxygen adsorbed at the current collector as well as the surface C-O species on the graphite4 are reduced resulting in resistivity drop which probably follows percolation mechanism. A careful analysis of the cyclic voltammetry curves measured at 5, 10, and 20 mV/s (Figure 10a-c) shows that the discharge capacity is approximately independent of ν whereas the Coulombic efficiency of the whole anodic process decreases as ν decreases. Thus the influence of the parasitic side reactions on the electrode’s resistivity is expected to be higher at relatively slow sweep rates, which is in accordance with our experimental finding.
J. Phys. Chem. B, Vol. 103, No. 9, 1999 1507 Conclusion The present study of the electrochemical oxidation of a composite graphite electrode in 4 M H2SO4 carried out by simultaneous application of various electroanalytical, optical, and structural sensitive techniques contributes to the understanding of the formation mechanism of graphite sulfates. Only dilute phase of graphite bisulfate (mainly VIII stage) can be obtained by polarization up to 1.5 V (vs silver wire in the same solution) which is in good agreement with the phase diagram for various graphite sulfates intercalation compounds proposed by Metrot et al.22 In situ XRD study shows that the dilute phase is formed continuously as a solid solution within the pristine graphite phase rather than a first-order phase transition reaction. The cyclic voltammetry curves obtained are in very good agreement with the dependence of the intensity of the crystallographic peak 002 of the pristine graphite phase with potential. The major parasitic reactions occurring at the electrode at high anodic potentials involves the oxidation of carbon atoms on the surface of graphite particles, as well as the adsorption and evolution of oxygen on graphite particles and the platinum current collector. Electronic resistance of the graphite composite electrode during its electrochemical cycling depends mainly on these layers formation and, to a lesser degree, on the variation in the intrinsic electronic in-plane resistivity of the host. The deflection of a probe laser beam near the electrodesolution interface toward the solution on oxidation and the opposite sign of the beam’s deflection on reduction provided direct evidence for the anion insertion-deinsertion to/from the graphite. The role of cointercalation of neutral H2SO4 and H2O species needs a special consideration, which should lead to a modification of the convolution procedure of the data for this particular case. Comparison between the electrochemical impedance characteristics of the graphite during its oxidation in H2SO4 and the reduction in a Li-containing aprotic electrolyte16-18 shows that in the latter case the electrode is stabilized by a resistive (and protective) solid-electrolyte interface (SEI). In contrast, during oxidation of graphite in aqueous H2SO4 SEI does not form, thus the side anodic reactions interfere with the anion insertion-deinsertion reaction. The chemical diffusion coefficient of HSO4- anion in the host matrix, estimated from the corresponding Nyquist plots, are 4-5 orders of magnitude larger than that for Li-ions. This may imply much higher attractive interactions between Li-ions and the graphite host as compared to those between HSO4- ions and the graphite. This is in good agreement with the well-known difference between the electronic structures of graphite’s donor and acceptor compounds: the former compounds experience a s-π-hybridization resulting in partial covalent bonding perpendicular to the graphene planes, whereas the latter compounds are characterized by a rather weak interactions between the intercalants and the carbon’s pz orbitals.27 Acknowledgment. This work has been partially supported by the French Ministry of Foreign Affairs and the APAPE (France) and the Ministry of Sciences and Arts (Israel) within the Arc-En-Ciel/Keshet collaboration program. We thank the French Embassy in Tel Aviv for their kind help in arranging the collaboration between the French and Israeli partners. We thank Prof. J. O. Besenhard of Gratz Technical University for the deep and stimulating discussions. References and Notes (1) Herold, A. In Intercalated Layered Materials; Levy, F., Ed.; D. Reidel Publishing Company: Dordrecht, Holland, 1979; p 355.
1508 J. Phys. Chem. B, Vol. 103, No. 9, 1999 (2) Besenhard, J. O.; Wudy, E.; Mohwald, H.; Nickl, J. J.; Bibiracher, W.; Foag, W. Synth. Met. 1983, 7, 185. (3) Aronson, S.; Lemont, S.; Weiner, J. Inorg. Chem. 1971, 10, 1296 (4) Metrot, A.; Fischer, J. E. Synth. Met. 1981, 3, 201. (5) Beck, F.; Krohn, H. Synth. Met. 1983, 7, 193. (6) Harrach, A.; Metrot, A. Electrochim. Acta 1989, 34, 1877. (7) Roger, J. P.; Fournier, D.; Boccara, A. C. J. Phys. 1983, 44, C6313. (8) Rosolen, J. M.; Facastoro-Decker, M.; Decker, F. J. Electroanal. Chem. 1993, 346, 119. (9) Weaver, J. K.; McLarnon, F. R.; Cairns, E. J. J. Electrochem. Soc. 1991, 138, 2579. (10) Vieil, E.; Meerholtz, K.; Matencio, T.; Heinze, J. J. Electroanal. Chem. 1994, 368, 183. (11) Barbero, C.; Miras, M. C.; Kotz, R.; Haas, O. Solid State Ionics 1993, 60, 167. (12) Orata, D.; Buttry, D. A. J. Am. Chem. Soc. 1987, 109, 3574. (13) Bruckenstein, A. R.; Wilde, C. P.; Shay, M.; Hillman, A. R. J. Phys. Chem. 1990, 94, 787. (14) Miras, M. C., Barbero, C.; Kotz, R.; Haas, O.; Schmidt, V. M. J. Electroanal. Chem. 1992, 338, 279.
Levi et al. (15) Vorotyntsev, M. A.; Vieil E.; Heinze, J. Electrochim. Acta 1996, 41, 1913. (16) Levi, M. D.; Aurbach, D. J. Phys. Chem. B, 1997, 101, 4630. (17) Levi, M. D.; Aurbach, D. J. Phys. Chem. B, 1997, 101, 4641. (18) Levi, M. D.; Levi, E. A.; Aurbach, D. J. Electroanal. Chem. 1997, 421, 89. (19) Mandelis, A.; Royce, B. S. H. Appl. Opt. 1984, 23, 2892. (20) Aurbach, D.; Levi, M. D.; Levi, E.; Teller, H.; Markovsky, B.; Salitra, G. J. Electrochem. Soc. 1998. In press. (21) Metrot, A.; Harrach, A. Electrochim. Acta 1993, 38, 2005. (22) Bouyad, B.; Marrouche, A.; Tihli, M.; Fuzellier, H.; Metrot, A. Synth. Met. 1983, 7, 159. (23) Ho, C.; Raistrick, I. D.; Huggins, R. A. J. Electrochem. Soc. 1980, 127, 343. (24) Fujii, R. Report of the GoVerment Industrial Research Institute 353, U.S. Government Printing Office: Washington, DC, 1978; p 1. (25) Metrot, A.; Tihli, M. Synth. Met. 1988, 23, 25. (26) Frumkin, A. N.; Fedorovich, N. V.; Kulakovskaya, S. I.; Levi, M. D. SoV. Electrochem. 1974, 10, 130. (27) Fischer, J. E. In Intercalated Layered Materials; Levy, F., Ed.; D. Reidel Publishing Company: Dordrecht, Holland, 1979; p 489.
