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In the format provided by the authors and unedited. © 2017 Macmillan Publishers Limited, part of Springer Nature. All rights reserved. SUPPLEMENTARY INFORMATION DOI: 10.1038/NMAT5029 NATURE MATERIALS | www.nature.com/naturematerials 1 A step toward safer and recyclable lithium-ion capacitors using sacrificial organic lithium salt Authors: P. Jeżowski 1 , O. Crosnier 2,3 , E. Deunf 2 , P. Poizot 2,4 , F. Béguin 1 , T. Brousse 2,3* 1. Poznan University of Technology, Institute of Chemistry and Technical Electrochemistry, ul. Berdychowo 4, 60-965 Poznan, Poland. 2. Institut des Matériaux Jean Rouxel, CNRS UMR 6502 – Université de Nantes, 2 rue de la Houssinière BP32229, 44322 Nantes Cedex 3 , France. 3. Réseau sur le Stockage Électrochimique de l’Énergie, FR CNRS 3459, 80039 Amiens Cedex, France. 4. Institut Universitaire de France (IUF), 1 rue Descartes, 75231 Paris Cedex 05, France * [email protected]
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In the format provided by the authors and unedited.

© 2017 Macmillan Publishers Limited, part of Springer Nature. All rights reserved.

SUPPLEMENTARY INFORMATIONDOI: 10.1038/NMAT5029

NATURE MATERIALS | www.nature.com/naturematerials 1

 

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Supplementary information

A step toward safer and recyclable lithium-ion capacitors using sacrificial organic lithium

salt

Authors:

P. Jeżowski1, O. Crosnier2,3, E. Deunf2, P. Poizot2,4, F. Béguin1, T. Brousse2,3*

1. Poznan University of Technology, Institute of Chemistry and Technical Electrochemistry, ul.

Berdychowo 4, 60-965 Poznan, Poland.

2. Institut des Matériaux Jean Rouxel, CNRS UMR 6502 – Université de Nantes, 2 rue de la

Houssinière BP32229, 44322 Nantes Cedex 3 , France.

3. Réseau sur le Stockage Électrochimique de l’Énergie, FR CNRS 3459, 80039 Amiens Cedex,France.

4. Institut Universitaire de France (IUF), 1 rue Descartes, 75231 Paris Cedex 05, France 

* [email protected]

© 2017 Macmillan Publishers Limited, part of Springer Nature. All rights reserved.

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S1. Characterisation of 3,4–dihydroxybenzonitrile dilithium salt (Li2DHBN)

Fourier Transform Infra-Red spectrometry (FTIR): a Bruker Vertex 70 device was used in the

wavenumber range from 4000 cm-1 to 400 cm-1. KBr pellets were prepared to record the KBr

spectra. Supplementary figure 1 shows the IR spectra for the KBr pellet of as-prepared Li2DHBN

after filtration (the dried Li2DHBN is extremely sensitive to moisture and reacted with the

atmosphere during the recording of its IR spectrum), and of 3,4–dihydroxybenzonitrile for

comparison purposes. Characteristic lines of tetrahydrofuran (THF) are easily identified in the

spectrum of as-prepared Li2DHBN (Supplementary figure 1b) in the range from 2882 to 2979

cm-1 (C–H stretching), and at 1041 cm-1 (C–O stretching). The proof of lithiation in figure S1 b is

evidenced by the disappearance of the characteristic O-H lines of 3,4–dihydroxybenzonitrile in

the range from ca. 3000 to 3500 cm-1 (Supplementary figure 1a). Interestingly, the spectrum of

Li2DHBN shows a characteristic band at around 488 cm-1, which, according to Zhu et al. [S1],

could be attributed to the Li–O vibration.

Supplementary figure 1 Infra-red spectra of a) 3,4–dihydroxybenzonitrile; and b) as-prepared

Li2DHBN.

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Differential Scanning Calorimetry (DSC) was performed with a DSC204 Phoenix 204 (Netzsch)

apparatus, where the protective gas was Nitrogen (N2) of analytical purity. The material (4 mg)

was placed in a closed aluminium crucible, and an empty crucible was used as a reference. The

samples were heated/cooled at 10 °C min-1 in the temperature range from 25 to 400 °C. Dried

Li2DHBN was introduced into the air-tight aluminium crucible in a glove box under argon

atmosphere and analysed by DSC (Supplementary figure 2). The absence of any signal at

temperatures lower than 100 °C confirms that THF (and water) were thoroughly eliminated

during drying. In the temperature range from 150 to 230 °C, there is slight deviation of the

baseline which is attributed to fast amorphisation, and is characteristic for lithiated organic

compounds. Between 285 and 305 °C, the Li2DHBN is exothermically decomposed with

formation of Li2CO3, as already observed for instance with tetrahydroxy-p-benzoquinone

tetralithium salt (Li4C6O6) [S2].

Supplementary figure 2 Differential scanning calorimetry thermogram of dried Li2DHBN at a

heating rate of 10 °C min-1.

Nuclear Magnetic Resonance (NMR) spectra were recorded with a Bruker AVANCE 500 MHz

spectrometer. Deuterated methanol (CD3OD) was used as the solvent and as the internal standard

for chemical shifts. In the 13C NMR spectrum of dried Li2DHBN (Supplementary figure 3), the

peaks at 165 ppm and 156 ppm are attributed to the –C–OLi carbons in the 3rd and 4th positions

of CN, respectively. The signals at 125 ppm and 117 ppm are related to the C–H carbons of the

aromatic ring. The shift at 124 ppm arises from the carbon atom of the CN group. The carbon

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atom of the aromatic ring close to the C≡N group produces the signal at 95 ppm. The peak at 50

ppm is assigned to the deuterated solvent (CD3OD).

Supplementary figure 3 13C Nuclear magnetic resonance spectrum of dried Li2DHBN at 500

MHz. Deuterated methanol (CD3OD) was used as the solvent.

The 1H nuclear magnetic resonance (NMR) spectrum of dried Li2DHBN is presented in

supplementary figure 4. The peak at 6.4 ppm, with an integration ratio of 1, characterises the

single aromatic proton in the ortho position of –CN. The multiplet at around 6.7 ppm, with an

intensity ratio of 2, is attributed to the two other aromatic protons. The peak at 3.2 ppm is due to

protons of the incompletely deuterated methanol. The peak at 4.8 ppm is characteristic of HDO

and is attributed to the partial decomposition of Li2DHBN into 3,4–dihydroxybenzonitrile due to

moisture present in the atmosphere or the deuterated methanol.

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Supplementary figure 4 1H nuclear magnetic resonance spectrum of dried Li2DHBN at

500 MHz. Deuterated methanol (CD3OD) was used as the solvent.

© 2017 Macmillan Publishers Limited, part of Springer Nature. All rights reserved.

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S2. Solubility measurement of Li2DHBN and electrochemical/thermal behaviours of both

the positive and negative electrodes in an LIC cell

Solubility evaluation of Li2DHBN in the relevant liquid electrolyte

The supplementary figure 5 shows visual evidence of the poor solubility of Li2DHBN in EC–

DMC/LiPF6 1 mol L-1 under real LIC cell conditions (i.e., 8 mg of Li2DHBN for 500 µL of

electrolyte).

Supplementary figure 5 Optical picture of powder of Li2DHBN in contact with EC-DMC/LiPF6

1 mol L-1.

The solubility of Li2DHBN in the electrolyte was then quantified by a simple spectrophotometric

titration method performed on diluted solutions (aromatic compounds being efficient absorbers of

UV radiation). The spectrophotometric data were recorded on a Cary UV-Vis spectrometer from

Agilent with a 1 cm-length quartz cell at = 290 nm (290 = 6809 L mol-1 cm-1). Experimentally,

a suspension of 8 mg of Li2DHBN with 500 µL of electrolyte (real LIC conditions) was prepared

and stirred for a few hours. After decantation and filtration, the surnatant was diluted 100 times

with the electrolyte solution prior to spectrophotometric measurements. The calibration curve

was performed between 0.01 to 0.08 mmol L-1 (supplementary figure 6) and the resulting

intrinsic solubility of Li2DHBN was found to be 0.95 g L-1 (= 950 ppm < 1%).

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Supplementary figure 6 UV-Vis calibration line together with the corresponding value defining

the intrinsic solubility of Li2DHBN in EC-DMC/LiPF6 1 mol L-1 (red dot).

Electrochemical behaviour

A free-standing positive electrode material with PTFE as the binder was prepared inside an

argon-filled glove box under the same conditions as described in the paragraph Methods of the

article. In order to estimate the oxidation limit of the positive electrode material, the anodic

polarisation of the composite electrode was prolonged to higher potentials. As evidenced in

supplementary figure 7, after the first irreversible lithium extraction plateau, there is a potential

rise followed by another endless irreversible plateau at a potential higher than 4.5 V vs. Li+/Li0,

which might be associated to electrolyte oxidation (Supplementary figure 7b). The exact value

of irreversible lithium extraction capacity was thus determined by the position of the inflexion

point, at ca. 360 mAh g-1 in supplementary figure 7a, thereby demonstrating that all the lithium

present in Li2DHBN (theoretical value of 365 mAh g-1) is available for graphite lithiation.

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Supplementary figure 7 Galvanostatic charge/discharge profiles of 3,4-dihydroxybenzonitrile

dilithium salt at C/10. The lithium extraction takes place until ca. 4.0 V vs. Li+/Li0. The second

plateau at ca. 4.5 V vs. Li+/Li0 is attributed to side reactions such as electrolyte oxidation. The

electrode was composed of 65 wt.% Li2DHBN, 30 wt.% carbon black and 5 wt.% PTFE binder.

The experiments were performed in a 1 mol L-1 LiPF6 dissolved in EC: DMC (vol. ratio 1:1)

electrolyte with metallic lithium as the counter/reference electrodes.

Operando spectroelectrochemical measurements of pure Li2DHBN were also performed using a

miniature fibre-optic spectrophotometer (FLAME-S-XR1-ES, Ocean Optics) in glove box. The

sacrificial organic salt was deposited by dip-coating onto an Indium tin oxide (ITO) electrode

using DMC as the dispersive solvent (supplementary figure 8 top), and then placed in a 1-cm-

long quartz cell filled with EC-DMC/LiPF6 1 mol L-1 (1.5 mL). The coated ITO electrode was

then oxidized to 4.4 V vs. Li+/Li0 using the Potentiostatic Intermittent Titration Technique (PITT)

(supplementary figure 9), and removed from the electrolyte. No deposit is to be seen on the ITO

electrode (supplementary figure 8 bottom), thus demonstrating the dissolution of the delithiated

form of Li2DHBN (i.e., DOBN) in the electrolyte. The presence of dissolved molecules was

further assessed by monitoring the UV-visible response (supplementary figure 10), which

clearly shows a change in the electrolyte spectrum after first oxidation of the Li2DHBN-coated

ITO electrode.

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Supplementary figure 8 Optical pictures of the ITO electrode (0.5 cm wide) covered by

Li2DHBN in the initial state (top) and after oxidation (bottom).

Supplementary figure 9 PITT charge profile of pure Li2DHBN deposited onto the ITO

electrode, measured during the in-operando spectroelectrochemical experiment in an argon-filled

glove box: 1-cm-long quartz cell filled with EC-DMC/LiPF6 1 mol L-1 (1.5 mL) and recorded

versus lithium metal.

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Supplementary figure 10 Corresponding UV-Vis response of the electrolyte after oxidation of

Li2DHBN deposited onto the ITO electrode to 4.4 V vs. Li+/Li0 by Potentiostatic Intermittent

Titration Technique. The concentration of DOBN is so great that the measured signal is saturated

in the UV region.

The same experiment was carried out in a Swagelok cell with a custom-made Teflon ring (11 mm

in diameter, 5 mm thick) where the amount of electrolyte was exactly 500 µL, which is the same

amount as in the separator of a full cell. After oxidation of the composite electrode (40 wt.% AC,

40 wt.% Li2DHBN, 15 wt.% Super C65 and 5wt.% PTFE) at 4.5V vs. Li+/Li0, the electrolyte was

removed from the cavity inside the Swagelok cell. As can be seen in supplementary figure 11,

the electrolyte turned a brownish colour.

Supplementary figure 11 Optical pictures of the electrolyte after oxidation of a composite

electrode (40 wt.% AC, 40 wt.% Li2DHBN, 15 wt.% Super C65 and 5wt.% of PTFE) at 4.5V vs

Li+/Li0

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The ionic conductivity of the electrolyte was measured after oxidation. No noticeable change was

observed between the ionic conductivity of the electrolyte in contact with the Li2DHBN-loaded

positive electrode, either at 293 K (10.4 ± 1.0 mS cm-1) or after oxidation (9.9 ± 1.0 mS cm-1).

Thus the oxidized form of Li2DHBN (i.e., DOBN), although completely dissolved in the

electrolyte, does not have significant influence on the ionic conductivity of the medium and

subsequently will not affect the power capability of our LIC.

Concomitantly, SEM images (supplementary figure 12) of electrodes, before a), b) and after

extraction of lithium from the organic molecule c), d), show that after the prelithiation process

using Li2DHBN (first charge) there are visible holes in surface of the electrode (one of the

regions with visible holes is marked with a red circle in supplementary figure 12c). This

observation is in agreement with the mass loss measured before and after lithium extraction

(solubilized DOBN molecules). Higher magnifications (b and d) show that the texture of the

electrode has changed after the lithium extraction process, including some holes and cracks,

which are not, however, detrimental to the cycling ability.

Supplementary figure 12 SEM images of composite positive electrodes (40 wt.% AC, 40

wt.% Li2DHBCN, 15 wt.% Super C65 and 5wt.% PTFE) before a), b) and after extraction of

lithium from the organic molecule Li2DHBN, which is the part of the positive composite

electrode c), d), at magnifications 100x a), c) and 200x b), d).

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Self-discharge measurements were performed on a two-electrode LIC cell, where the positive

electrode was loaded with Li2DHBN. A lithium reference electrode was added to monitor the

behaviour of individual electrodes (supplementary figure 13). After the first oxidation of the

composite positive electrode, the cell was polarized at 4 V for 2 hours, and then the open-circuit

voltage and potential were recorded.

Supplementary figure 13: Self-discharge plot of an LIC with a sacrificial composite positive

electrode based on Li2DHBCN. The composition of the positive electrode is 40 wt.% AC, 40

wt.% Li2DHBCN, 15 wt.% Super C65 and 5wt.% PTFE. The electrolyte is 1 mol L-1 LiPF6 in

EC: DMC. The black solid line represents the cell voltage; the red and the blue lines are the

potential profiles of the positive and negative electrodes, respectively.

According to supplementary figure 13, the positive electrode is mainly responsible for the

voltage decrease during the 20-hour OCV period. The potential fade of the positive electrode is

equal to 16% after 20 hours, i.e. it drops from 4.0 V vs. ref. Li+/Li0 to 3.6 V vs. ref. Li+/Li0.

Correspondingly, the potential of the negative electrode increased from 90 mV vs. ref. Li+/Li0 to

135 mV vs. ref. Li+/Li0 during the same OCV period. From the supplementary figure 13, it is

clear that there is no redox shuttle effect due to the dissolved DOBN molecules. Indeed, a 0.4V

loss over 20 hours is a commonly measured value for lithium-ion capacitors with an LP30

electrolyte [S3], and for symmetrical activated carbon electrochemical capacitors [S4].

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Thermal behaviour of the Li2DHBN/AC composite positive electrode in an LIC cell

DSC measurements of both the positive and the negative electrodes of LIC cells were performed

before (OCP state) and after the first charge (supplementary figure 14). These experiments were

conducted using a Q20 DSC (TA instruments) heat-flux differential calorimeter at a heating rate

of 10 K/min in a temperature range of 25–400 °C under a constant argon flow of 200 mL/min.

Experimentally, the LIC cells were carefully dismantled in a glove box and the recovered

samples (studied material with its electrolyte) were introduced into an aluminium crucible, which

was then sealed. The loaded DSC crucibles were pierced prior to measurement.

Supplementary figure 14 Galvanostatic profile upon the first charge of an LIC cell using 66

wt% Li2DHBN mixed with 33 wt% carbon black (Super C65, Imerys) as the composite positive

electrode, and graphite (SLC 1512P, Superior Graphite) as the negative electrode, recorded at a

cycling rate of C/10 using EC–DMC/LiPF6 1 mol L-1 as the electrolyte.

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Supplementary figure 15 Typical DSC traces of the positive electrode (66 wt% Li2DHBN and

33wt% carbon black) and negative electrode (graphite), gathered a) before cycling the cell and b)

after the first charge, without rinsing the electrodes after battery dismantling. All DSC

measurements were recorded under Ar flow (200 mL min-1) in a pierced crucible.

After the first charge, both electrodes exhibited a weak exothermic peak at 120°C and 150°C for

the positive and negative electrodes, respectively (supplementary figure 15b). The heat

generated at the charged positive electrode soaked in the electrolyte was a mere 63 J g-1. That at

the charged negative electrode is in the same order of magnitude (79 J g-1). These values are well

below those reported in literature for lithiated graphite electrodes in lithium-ion batteries using

different graphite electrodes, where it was found that the exothermic peak corresponds to 725 J g-

1 [S5].

S3. Alternative prelithiation conditions of the graphite negative electrode

Supplementary figure 16 shows the electrochemical prelithiation of graphite achieved by

galvanostatic cycling, where the S.E.I. was formed at C, and intercalation at C/10 (C is the

theoretical capacity of graphite, i.e. 372 mAh g-1). The potential limits were fixed at 4.00 vs.

Li+/Li0 for the positive electrode and at 0.01 V vs. Li+/Li0 for the negative one, in order to prevent

electrolyte oxidation and lithium plating, respectively. When the potential of the negative

electrode reached 0.20 V vs. Li+/Li0 (at which point it can be considered that the S.E.I. is totally

formed), the potential of the positive electrode was ca. 3.95 V vs. Li+/Li0, which is very close to

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the positive electrode limit. Therefore, at this point, the cell was left at open circuit potential for 2

h, after which the lithium was intercalated at C/10. These conditions were already optimised

through our former research dedicated to the use of Li5ReO6 as the irreversible lithium source for

LICs [S6]. In supplementary figure 16 it can be seen that the total capacity is around 250 mAh

g-1, which is much lower than the expected 360 mAh g-1 found in figure 3b, when the current is

fixed at C/10. Hence, it seems that the presence of Li2DHBN modifies the optimal conditions of

the prelithiation step. This thus implies that a certain amount of Li2DHBN might have been

dissolved, probably due to the use of an excessively high current (C) during the S.E.I. formation.

Supplementary figure 16 Profiles for the electrodes potential and cell voltage during

galvanostatic S.E.I. formation and graphite lithiation facilitated by a composite positive

electrode composed of 40 wt.% AC, 40 wt.% Li2DHBN, 15 wt.% carbon black and 5 wt.%

PTFE. The S.E.I. was formed at C until reaching a potential of 0.2 V vs. Li+/Li0 (area highlighted

in green); then, after a rest period of 2 h, the lithium was intercalated into the graphite at C/10

(area highlighted in yellow). The potential limits for the positive and negative electrodes are 4.0

V and 0.01 V vs. Li+/Li0, respectively. Colour of the curves: dashed red for the potential of the

positive electrode; dotted blue for the potential of the negative electrode; black for the cell

voltage profile. The electrolyte is 1 mol L-1 LiPF6 in EC: DMC.

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S4. Optimisation of cell voltage range for an LIC cell obtained after graphite prelithiation

at C/10

In order to enhance the energy of the LIC based on the renewable lithium source, we tried

slightly increasing the maximum voltage value to 4.1 V. The prelithiation was performed under

the conditions shown in figure 3b. Unfortunately, as shown in supplementary figure 17, there is

a slight but visible continuous capacitance decay, which might be connected to electrolyte

oxidative decomposition, and this decay accelerates when the current is set to 0.65 A g-1 after

1350 cycles. Hence, to promote the stable operation of the LIC based on a positive electrode

containing delithiated Li2DHBN, it is necessary to restrict the upper voltage limit to 4 V, as

confirmed by supplementary figure 17. This value seems to be even slightly higher than that of

3.8 V put forward by JM Energy for the ULTIMO system [S7].

Supplementary figure 17 Cycle life of an LIC with a sacrificial composite positive electrode

based on Li2DHBN where S.E.I.. formation and prelithiation was performed at C/10. The LIC

was cycled at 0.25 A g-1 (blue squares), 0.50 A g-1 (green circles) and 0.65 A g-1 (red diamonds)

in the voltage range from 2.2 to 4.1 V. The electrolyte was 1 mol L-1 LiPF6 in EC:DMC.

In order to avoid possible plating at higher currents, the upper voltage limit should be decreased,

which can be achieved by using a system such as the ULTIMO (constructed by JM Energy) that

is operated within a 2.2 to 3.8 V voltage range. Three separate two-electrode cells with reference

electrode were assembled and tested in different voltage ranges seen in supplementary figure 18

a), b) 2.2 – 4.0 V, c), d) 2.2 – 3.9 V and e), f) 2.2 – 3.8 V at current 0.5 A g-1. The plots a), c) and

e) present the initial galvanostatic charge/discharge cycles, whereas plots b), d), f) show those

recorded after 1000 cycles at corresponding voltage ranges. A noticeable difference was observed

in the lowest values of potential for the negative electrode: 23 mV vs. ref. Li+/Li0, 29 mV vs. ref.

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Li+/Li0and 33 mV vs. ref. Li+/Li0 when the systems reached their maximum voltage of 4.0 V, 3.9

and 3.8 V, respectively. This serves to demonstrate that it is possible to guard the LIC system

against plating by adjusting the maximum voltage. The minimum potential of the negative

electrode after 1000 cycles was 24 mV vs. ref. Li+/Li0, 30mV vs. ref. Li+/Li0 and 34 mV vs. ref.

Li+/Li0 for the respective voltages of 4.0 V, 3.9 V and 3.8 V.

Supplementary figure 18 Galvanostatic charge/discharge profiles of a two-electrode LIC with

reference electrode and with sacrificial composite positive electrode based on Li2DHBCN at 0.50

A g-1 (current per gram of graphite) in the voltage range of 2.2 V to 4.0 V a), b); to 3.9 V c), d);

to 3.8 V e), f). The composition of the positive electrode is 40 wt.% AC, 40 wt.% Li2DHBCN, 15

wt.% Super C65 and 5wt.% PTFE. The electrolyte is 1 mol L-1 LiPF6 in EC: DMC. The black

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solid line represents the cell voltage; the red and the blue lines are the potential profiles of the

positive and negative electrodes, respectively. For the purpose of monitoring the potential of each

electrode, it was necessary to use a1550-μm-thick separator.

References

[S1] Cai, W., Wang, H., Sun, D., Zhang, Q., Yao, X. & Zhu, M. Destabilization of LiBH4

dehydrogenation through H+–H- interactions by cooperating with alkali metal hydroxides. RSC

Adv. 4, 3082–3089 (2014).

[S2] Chen, H., Armand, M., Courty, M., Jiang, M., Grey, C.P., Dolhem, F., Tarascon, J.-M. &

Poizot, P. Lithium salt of tetrahydroxybenzoquinone: toward the development of a sustainable Li-

ion battery. J. Am. Chem. Soc. 131, 8984–8988 (2009).

[S3] Decaux, C., Lota, G., Raymundo-Piñero, E., Frackowiak, E. & Béguin F. Electrochemical

performance of a hybrid lithium-ion capacitor with a graphite anode preloaded from lithium

bis(trifluoromethane)sulfonimide-based electrolyte, Electrochim. Acta 86, 282–286 (2012).

[S4] Hao, C., Wang, X., Yin, Y., & You, Z. Analysis of Charge Redistribution During Self-

discharge of Double-Layer Supercapacitors, J. Electron. Mater. 45, 2160–2171 (2016).

[S5] Eshetu, G., Grugeon, S., Gachot, G., Mathiron, D., Armand, M., & Laruelle, S. LiFSI vs.

LiPF6 electrolytes in contact with lithiated graphite: Comparing thermal stabilities and

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[S6] Jeżowski, P., Fic, K., Crosnier, O., Brousse, T. & Béguin, F. Lithium rhenium(VII) oxide as

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[S7] http://www.jmenergy.co.jp/en/, accessed January 2017.


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