Aqueous Intercalation of Graphite at a Near-Neutral pH - ACS Applied

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Aqueous Intercalation of Graphite at a Near-Neutral pH Francis M Alcorn, Kaci L Kuntz, Daniel L. Druffel, and Scott C. Warren ACS Appl. Energy Mater., Just Accepted Manuscript • DOI: 10.1021/acsaem.8b01101 • Publication Date (Web): 21 Aug 2018 Downloaded from http://pubs.acs.org on August 27, 2018

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ACS Applied Energy Materials

Aqueous Intercalation of Graphite at a Near-Neutral pH

Francis M. Alcorn§, Kaci L. Kuntz§, Daniel L. Druffel§, and Scott C. Warren§†*

§

Department of Chemistry and †Department of Applied Physical Sciences, University of North

Carolina at Chapel Hill, Chapel Hill, NC 27599, USA. * [email protected]

Keywords: GIC, aqueous battery, in-situ x-ray diffraction, intercalation chemistry, anion intercalation, exfoliation

ACS Paragon Plus Environment

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ABSTRACT:

Graphite has been extensively studied as a battery electrode, but few investigations have explored its behavior in water at neutral pH. Here, we investigate graphite intercalation in an aqueous ammonium sulfate electrolyte with a pH of 6.

We identified potentials at which

bisulfate intercalation and deintercalation compete with water oxidation.

In-situ X-ray

diffraction revealed that the interlayer space of graphite expands during intercalation, reaching a stage-two graphite intercalation compound. Irreversible changes to the graphite occurred during intercalation, as demonstrated by the appearance of sp3-type carbon, hydroxyl, and carbonyl groups in Raman and X-ray photoemission spectra. Our findings indicate that intercalation of graphite in aqueous media at a near-neutral pH and modest potentials is possible, but improved chemical and structural stability of the graphite and electrolyte are required to achieve long-term operation.


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Batteries with aqueous electrolytes could enable inexpensive grid-scale energy storage.1-7 Because water facilitates many undesired electrochemical reactions, such as electrode oxidation or water splitting, it remains a challenge to develop suitable electrodes and electrolytes.8 Ionintercalated graphite has not received much attention for neutral-pH aqueous batteries, even though it is widely used in battery electrodes.9,10 Most graphite intercalation compounds (GICs) are unstable in water,11-13 and the few ions that have been intercalated in aqueous electrolytes only do so at extreme conditions, such as pH < 0 (e.g., 10 M sulfuric acid)14-21. As a result, to the best of our knowledge, ions have never been intercalated into graphite at neutral pH in more than trace quantities,12 and relatively little is known about the chemical or physical behavior of near-neutral pH intercalation electrodes, including GICs. The window of water’s electrochemical stability naturally imposes limitations on the type of ions that can be intercalated into graphite. Undoped graphite’s Fermi level, 4.7 eV from vacuum,22 is closer to water’s reduction potential (0.2 eV above graphite’s Fermi level) than water’s oxidation potential (1.0 eV below graphite’s Fermi level). We therefore anticipated that water decomposition reactions would be minimized if anions were intercalated rather than cations, since we could apply positive potentials up to ca. +1 V for anion intercalation before water decomposition would become thermodynamically favorable.

Moreover, the large

overpotentials for water oxidation23 could allow potentials more than +1 V to be applied to graphite without significant water splitting reactions. This strategy has succeeded in a different aqueous system, the lead acid battery, where the positive electrode operates at ca. 1.6 V vs. NHE.24,25

Of the many possible anions, we explored the sulfate intercalation system

(SO42-/HSO4-) because sulfuric acid/graphite is the most well-studied aqueous GIC14-19,26, albeit at pHs < 0. We selected an aqueous ammonium sulfate electrolyte, which has a pH of 5.8, to


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investigate the intercalation of graphite at a near-neutral pH. Our investigation of this system was inspired by a recent study27 that showed an ammonium sulfate electrolyte could intercalate graphite and subsequently exfoliate it to produce graphene. It is notable, however, that the authors applied a potential of +10 V to the graphite, which facilitated exfoliation but prevented isolation or characterization of a GIC. We therefore build on this previous study to explore the sulfate system for graphite intercalation at modest potentials. Here we study reversible electrochemical intercalation of graphite in a 2 M (NH4)2SO4 electrolyte using several voltammetric and potentiometric techniques, as well as in-situ X-ray diffraction (XRD). In-situ XRD indicates that intercalation does occur with the formation of well-staged GICs up to 2. The stage number describes the number of consecutive graphene layers between intercalant layers.28 In-situ Raman spectroscopy and ex-situ X-ray photoemission spectroscopy (XPS) show, however, that intercalation is accompanied by oxidative damage of the material. Although this damage limits the reversibility of the graphite/sulfate system, our work demonstrates the first staged GIC produced via intercalation in near-neutral pH water. Graphite electrodes were prepared by drop-casting 90:10 (w:w) graphite:poly(vinylidene fluoride) (PVDF), where PVDF is used as a binder. We compared the electrochemical behavior of a 2 M ammonium sulfate electrolyte, pH 5.8, and a non-intercalating 2 M ammonium phosphate ((NH4)1.33H1.67PO4) buffer, pH 6.1.17,29 We used a graphite working electrode, platinum counter electrode, and Ag/AgCl reference electrode. For in-situ XRD and Raman spectroscopy, graphite films were deposited onto platinum in a planar two-electrode electrochemical cell (Figure S1) with a platinum counter electrode. For XRD, Raman spectroscopy, and XPS studies, a constant current (2-4 mA/cm2) was applied to drive intercalation.


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Figure 1. (A) Linear sweep voltammetry (LSV) from 1.6 to 1.95 V vs. Ag/AgCl (scan rate = 0.150 mV/s) for graphite/PVDF (solid lines) or glassy carbon (dashed lines) electrodes in ammonium sulfate or ammonium phosphate electrolytes. (B) Chronopotentiograms for the constant current oxidation of graphite in both electrolytes (3 mA/cm2 for sulfate, 4 mA/cm2 for phosphate), and (C) chronoamperometry at 1.6 V vs. Ag/AgCl following the LSV in part (A). In all plots, red lines denote a phosphate electrolyte and blue lines denote sulfate. To study the electrochemical behavior of the graphite-sulfate system, linear sweep voltammetry (LSV) from 1.6 to 1.95 V vs. Ag/AgCl (Figure 1A) was used to drive oxidation at a graphite working electrode in both an ammonium sulfate and a non-intercalating17,29 ammonium phosphate electrolyte. The sulfate electrolyte showed a maximum current at 1.8 to 1.9 V vs. Ag/AgCl and a markedly larger current density than the phosphate electrolyte. Equivalent scan rates were used in these experiments (0.150 mV/s), so the additional current in the sulfate system is primarily due to a Faradaic process rather than a capacitive one. Furthermore, a representative chronopotentiogram (Figure 1B) in the sulfate electrolyte displayed a voltage plateau around 1.9 V vs. Ag/AgCl, in agreement with the peak voltage in LSV (Figure 1A), suggesting that a steady-state Faradaic process is occurring, such as intercalation.

In contrast, a

chronopotentiogram for the system in the phosphate electrolyte (Figure 1B) shows a gradually


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increasing voltage above 2.0 V vs. Ag/AgCl.

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To determine whether electrolysis of water

accounted for the observed differences between sulfate and phosphate electrolytes, LSV was repeated using a glassy carbon working electrode, where intercalation reactions are not possible (Figure 1A, dashed lines). The water splitting currents were similar in both electrolytes and thus could not account for the differences observed in the graphite electrode (Figure 1A, solid lines). This suggests that oxidative intercalation of graphite or oxidative damage of graphite is more favorable in the sulfate electrolyte than in the phosphate electrolyte. The observation of negative current density at 1.6 V vs. Ag/AgCl (Figure 1C) in the sulfate electrolyte immediately following the LSV experiment in Figure 1A is consistent with a deintercalation process, which implies that the Faradaic current observed in the sulfate electrolyte is at least partially due to sulfate intercalation. The charge passed during deintercalation was smaller than during intercalation, indicating that the intercalation was not fully reversible. We calculate that 15 ± 10% (N = 4) of the intercalant successfully deintercalates.

We hypothesize that the low deintercalation

efficiency is due to damage to the graphite, which we discuss further below. The intercalation potential of graphite in 2 M ammonium sulfate (1.8 to 1.9 V vs. Ag/AgCl) is higher than the corresponding potential in concentrated sulfuric acid (as low as 0.9 V vs. Ag/AgCl).14 This difference is perhaps unsurprising given the significantly lower concentration of intercalant in the ammonium sulfate electrolyte. The interlayer chemistry of the sulfate system is complicated, however, and can include water, S2O82-, and hydronium, in addition to bisulfate and sulfuric acid.30,31 Consequently, because the identity of the intercalant(s) may be complex, it is difficult to determine whether the roughly +1 V shift in intercalation potential from sulfuric acid to ammonium sulfate can be explained solely from the lower concentration of intercalant(s).


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Figure 2. (A) In-situ XRD of graphite intercalated in 2 M ammonium sulfate. The emergence of the GIC peaks ({00n} and {00n+1}) and the reduction of intensity of the graphite {002} peak are consistent with intercalation. Times shown denote the time of intercalation at the start of each scan. The Si {111} intensity decreases as the mass of the graphite electrode increases with intercalation. (B)-(F) Illustration of planes that give rise to diffraction in graphite (B) and various GICs (C)-(F). Orange = intercalant layer; gray = graphene layer. There is 3.35 Å between graphene layers and 2.1 Å between graphene and intercalant layers. To obtain insight into the structural and chemical changes that accompanied the electrochemical oxidation, we employed in-situ XRD. XRD patterns (Figure 2A) were collected at 12-minute intervals at 2-ϴ angles from 23° to 32° during the constant current oxidation (j = 2 mA/cm2) of graphite in a (NH4)2SO4 electrolyte using a custom two-electrode cell (Figure S1). Si powder was added to the sample to provide an internal standard at 2-ϴ = 27.5°. Initially, only one peak from the graphite was present (2-ϴ = 26.6°), corresponding to graphite’s {002} family


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of planes with the expected interlayer separation of 3.35 Å (Fig. 2B). Within 13 minutes of constant-current oxidation, the {002} intensity decreased and a new peak at 25.8° emerged. Since the new peak was at a lower angle than the initial graphite peak, it suggested an expansion of the interlayer distance and is consistent with the formation of a GIC. We assigned the new peak to the {00n} set of planes32 of the GIC, with n denoting the stage number. Continued intercalation shifted the 00n peak to lower angles, indicating further interlayer expansion. By 37 min. of intercalation, a broad, low intensity peak emerged at ca. 30°, which has been seen previously in high-stage GICs, and we label as {00n+1}.32-33 The origin of the {00n} and {00n+1} diffraction in GICs is the following: {00n} arises from constructive interference from planes of graphene alone, while {00n+1} arises from planes of graphene and planes of intercalant (Fig. 2C-F). Because intercalant expands the distance between graphene planes, {00n} shifts to lower 2-ϴ. Meanwhile, because the distance between a plane of intercalant and an adjacent plane of graphene is smaller than the interlayer separation in graphite, the {00n+1} shifts to higher 2-ϴ. From the positions of these peaks, the stage of the high stage-GIC was calculated from the following equation32: =

1 sin −1 sin

1 .

Based on equation 1, we calculated that the GIC peak at 2-ϴ = 25.4° at 110 min. corresponds to n = 4.97, which we label hereafter as a stage 5 GIC. The observation of stage 5 at 25.4° agrees with a prior study of a stage 5 sulfuric acid GIC.14 At longer intercalation times, additional peaks emerged at 24.4° and 23.6°, in agreement with the reported peak positions of stage 3 and stage 2 sulfuric acid GIC, respectively.14 We did not observe a stage 4 GIC by XRD.


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Interestingly, the average interlayer spacing of the stage n GIC decreased upon appearance of the stage 3 GIC (see 00n peak in Figure 2 from 50 min. to 73 min.). We attribute this to the conversion of a stage 6 GIC (6 graphite layers between intercalant layers) into stage 3 GIC (3 graphite layers between intercalant layers), leaving the remaining 00n GIC with greater n > 6 character. The direct conversion of stage 6 into stage 3 may explain why stage 4 was not observed. From the peaks for stages 2, 3, and 5 (Figure 2), the interlayer space of the intercalated layer— known as the gallery height—was found to be 4.2 Å. This is identical to the previous report of graphite intercalated in concentrated sulfuric acid, where the intercalating species is the bisulfate ion (HSO4-).14 However, both sulfuric acid and the sulfate ion are similar in size to bisulfate, therefore the gallery height alone does not allow us to identify the intercalant(s) in this system. In fact, bisulfate is unlikely to be the intercalant, because the concentration of bisulfate in a 2 M ammonium sulfate electrolyte at pH 5.8 is just 0.3 mM (the pKa of bisulfate is 2.0). It is more likely that the intercalant is sulfate or a sulfate derivative, which we discuss further below.

Figure 3. In-situ Raman spectroscopy (λ0= 633 nm) of graphite (A) before intercalation and (B) after 30 min. of current-driven intercalation. Appearance of a shoulder on the G peak in B is


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consistent with intercalation of an acceptor-type species, and the emergence of the D peak indicates damage to the material.17 The G peak was fit to two peaks centered at 1611 cm-1 (blue) and 1583 cm-1 (red), corresponding to graphite sheets adjacent and non-adjacent to intercalant layers, respectively. Our observation of a stage 2 GIC (Figure 2) in a near neutral pH aqueous electrolyte is in stark contrast to prior reports that have suggested that concentrated acids are needed for intercalation.18,19 To confirm our observation and better understand the corresponding chemical and electronic changes, we employed in-situ Raman spectroscopy. Using a two electrode electrochemical cell (Figure S1), we collected Raman spectra of a graphite electrode before (Figure 3A) and during (Figure 3B) intercalation. Initially, the film was largely free of defects, as there was only one peak present in the Raman spectrum at 1582 cm-1, which is attributed to the carbon G peak, denoting sp2 carbon (Figure 3B). After 30 min. of intercalation (Figure 3B), a broad shoulder appeared on the G peak and was centered at 1611 cm-1, consistent with the intercalation of an acceptor-type species such as an anion.17,34 The position of this shoulder corresponds approximately to a stage 3 GIC,17 although the area of the 1611 cm-1 peak is three times lower and the peak is broader than expected for stage 3.17 These observations suggest that after 30 minutes, there are a mixture of stages, consistent with XRD. During intercalation, we also observed the emergence of a peak at 1360 cm-1. This has previously been assigned as the D peak of graphite and is attributed to the formation of sp3 carbon.

The formation of sp3 defects in GICs was reported to occur at sulfuric acid

concentrations below 2 M,17 likely from carbon oxy/hydroxide species as the GIC reacts with water.13 Collectively, our observations and prior studies may suggest that damage to graphite would be minimized by lowering the concentration of water and raising the concentration of


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electrolyte—a strategy that is now being widely investigated in multiple types of aqueous batteries.8,35-37

Figure 4. XPS of graphite electrodes and GICs after Ar-ion etching. (A) The carbon 1s spectra before (top) and after (below) intercalation. Peaks were fit to several carbon species: graphitic (orange), PVDF and hydroxyl (blue), fluorinated (PVDF, green), and carbonyl (magenta). The (B) sulfur 2p and (C) nitrogen 1s spectra are shown before (blue) and after (red) intercalation. The sulfur 2p peak was fit with a doublet for the 2p3/2 and 2p1/2 electrons. To better identify the defects that formed during intercalation, we employed XPS. Prior to the XPS measurement, the samples were washed in DI water and sputtered with Ar-ions to remove surface adsorbates. Spectra for the graphite electrodes before and after intercalation are presented in Figure 4A.. The graphitic carbon 1s peak (284.8 eV) broadened following intercalation (Table S1), consistent with damage to the graphite lattice.38,39 Before intercalation, a peak at 286.0 eV was observed, consistent with the C-H bonds of the PVDF binder, along with a peak at 290.9 eV,


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assigned to the C-F bonds of PVDF.40 After intercalation, C 1s intensity increased near 286 eV, which we attributed to the formation of a graphitic C-OH (Table S1).41-43 In addition, a small peak was observed at 288.6 eV after intercalation, and has been previously identified as a carbonyl group.41 Thus, in agreement with Raman spectra, XPS reveals oxidative damage of the graphite. We concluded that this is principally due to the formation of hydroxyl groups. We also investigated the sulfur 2p (Figure 4B) and nitrogen 1s (Figure 4C) signals by XPS to identify the possible presence of intercalated species.30 As noted above, we sputtered the sample with Ar ions to probe the sub-surface composition.

Before intercalation, there was no

identifiable signal from either N or S, but after intercalation, strong N and S signals were detected, corresponding to a N:S mole ratio of 58:42.

The nitrogen binding energy was

consistent with NH4+,44 while the S was consistent with either SO42- or HSO4-.45 The nearly 1:1 ratio of N:S, along with the Raman spectra that identified the intercalant as an acceptor, allows us to propose that the intercalant is the NH4SO4- anion. This proposal is consistent with the gallery height of 4.2 Å, a lattice spacing that is consistent with the HSO4- anion.14 In either the NH4SO4- or HSO4-, the gallery height is presumably determined by the diameter of the SO4 tetrahedron. This would imply that the NH4+ or H+ are co-planar with the sulfate, stabilizing the intercalant layer through hydrogen bonding.46-48


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Figure 5. (A) LSV from 1.6 to 1.95 V vs. Ag/AgCl (scan rate 0.150 mV/s) for graphite/PVDF in ammonium sulfate (light blue), magnesium sulfate (pH 6.4, purple), sodium sulfate (pH 5.7, green), and ammonium phosphate (pH 6.1 ± 0.1, red) electrolytes (2 M sulfate or phosphate), followed by (B) chronoamperometry at 1.6 V vs. Ag/AgCl in each electrolyte. To test our proposal that the ammonium cation co-intercalates with sulfate, we examined the electrochemical behavior of the graphite/PVDF electrode in 2 M magnesium sulfate (pH = 6.4) and 2 M sodium sulfate (pH = 5.7). We compare the behavior to that of the electrode in 2 M ammonium sulfate (pH = 5.8) and the non-intercalating phosphate electrolyte (pH = 6.1) in Figure 5. In LSV from 1.6 to 1.95 V vs. Ag/AgCl (Figure 1A), current densities were noticeably different between the three sulfate electrolytes, and all were larger than the phosphate electrolyte. Immediately following the potential sweeps to 1.95 V vs. Ag/AgCl, we lowered the applied potential to 1.6 V vs. Ag/AgCl and measured the resulting current (Figure 5B). Each of the sulfate electrolytes showed a non-zero current. From the data in Figure 5A and 5B, we conclude that intercalation is occurring in each of the sulfate electrolytes. The differences in current density seen in Figure 5A are consistent with the co-intercalation of a cation: if the intercalant species was only SO42- or HSO4-, the current densities would not differ significantly. In the case


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of the sodium sulfate and magnesium sulfate electrolytes, the intercalants could be NaSO4- and Mg(SO4)22-, but further analysis is required to confirm this proposal. In summary, we demonstrated the intercalation of the graphite up to a stage 2 GIC in an aqueous electrolyte at near-neutral pH. This is in marked contrast to all prior studies of aqueous graphite intercalation, to the best of our knowledge, which have been performed in concentrated acids.18,19 We observed oxidative intercalation of the graphite by an anion in addition to oxidative damage to the graphite itself, as shown by the emergence of a D peak in the Raman spectra during intercalation17 and changes of the carbon 1s XPS signal.38,39,41 Addressing graphite’s stability will be required for the development of viable graphite-based aqueous battery materials. An ion that more favorably intercalates into graphite could be of interest, as this would likely lower the potential required to drive intercalation. Alternatively, applying the water-in saltstrategy8,35-37 to a sulfate-based electrolyte may improve graphite intercalation, as the decreased water activity would limit the reactions with the GIC. This work suggests that reversible intercalation of GICs in near-neutral pH electrolytes for battery applications may be achievable.

ASSOCIATED CONTENT Supporting Information. Experimental details, Additional XPS data. AUTHOR INFORMATION email address: [email protected]. URL of the group website: http://warren.chem.unc.edu Notes The authors declare no competing financial interest.


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ACKNOWLEDGMENTS S.C.W. acknowledges funding support from the Beckman Young Investigator Award. F.M.A. acknowledges funding support from Scotty and Sonny Jackson and the UNC Department of Chemistry for the Matthew Neely Jackson Summer Undergraduate Research Award. This work was performed in part at the Chapel Hill Analytical and Nanofabrication Laboratory, CHANL, a member of the North Carolina Research Triangle Nanotechnology Network, RTNN, which is supported by the National Science Foundation, Grant ECCS-1542015, as part of the National Nanotechnology Coordinated Infrastructure, NNCI. This work made use Renishaw inVia Raman microscope in the UNC EFRC Instrumentation Facility established by the UNC EFRC: Center for Solar Fuels, an Energy Frontier Research Center funded by the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences under Award Number DE-SC0001011.

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Aqueous Intercalation of Graphite at a Near-Neutral pH - ACS Applied

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