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The Li–S battery: an investigation of redox shuttle
and self-discharge behaviour with LiNO3containing electrolytes†
Matthew J. Lacey,*a Anurag Yalamanchili,ab Julia Maibach,a Carl Tengstedt,bKristina Edströma and Daniel Brandella
The polysulﬁde redox shuttle and self-discharge behaviour of lithium–sulfur (Li–S) cells containing theelectrolyte additive LiNO3 has been thoroughly explored by a range of electrochemical and surface

Received 9th November 2015
Accepted 12th December 2015

analysis techniques on simple Li–S (i.e., not speciﬁcally optimised to resist self-discharge) andsymmetrical Li–Li cells. Despite the relatively eﬀective passivation of the negative electrode by LiNO3,
fully charged cells self-discharged a quarter of their capacity within 3 days, although in the short-term
cells can be recharged without any noticeable capacity loss. The processes governing the rate and

DOI: 10.1039/c5ra23635e

reversibility of self-discharge in these cells have been investigated and explained in terms of the

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reactions of polysulﬁdes occurring at both electrodes during idle conditions.

1

Introduction

The rechargeable lithium–sulfur (Li–S) battery is currently oneof the most actively studied “post-Li-ion” energy storage
systems, primarily because of its high energy density: practicalenergy densities on the cell level in excess of 300 W h kg1 –higher than the state of the art Li-ion batteries – have been
demonstrated in the private sector and energies in the range
400–600 W h kg1 are widely considered to be practicallyachievable. Other attractive advantages include the low cost of
the active material and good safety, although intrinsic drawbacks of poor conductivity of the active materials and severe
issues of parasitic reactions resulting from the solubility ofreaction intermediates (polysuldes, oen referred to as Li2Sn
or simply “PS”) are a barrier to wider practical application.
Signicant advances have been made in recent years in thedevelopment of the positive electrode in this system; these havebeen covered by numerous recent reviews and it is not necessary
to discuss these here.1–4 However, management of the behaviourof dissolved polysuldes remains a topic of considerable
interest and importance with a broad range of strategies underconsideration.
Polysuldes are readily soluble in most of the commonly
investigated electrolyte systems. The unwanted spontaneous
reaction of Li2Sn with the reactive Li metal negative electrode is

Department of Chemistry – Ångström Laboratory, Box 538, Lägerhyddsvägen 1, 751
21 Uppsala, Sweden. E-mail: matthew.lacey@kemi.uu.seab

Scania CV AB, SE-151 87 Södertälje, Sweden† Electronic supplementary information (ESI) available: Reference XPS spectraand supporting plots for redox shuttle rate determination. See DOI:
10.1039/c5ra23635e

3632 | RSC Adv., 2016, 6, 3632–3641

the source of the well-known “polysulde redox shuttle”,
eﬀectively an internal shuttle for electrons carried by multiple
polysulde species.
This problem could logically be tackled by designing theelectrolyte to minimise the solubility of Li2Sn species in theelectrolyte, as has been variously reported.5–9 However, thesolubility and reactivity of these species are important
contributors to the overall cell reaction10,11 – so suppression of
their solubility likely comes at the cost of decreased energy
eﬃciency and power density. A more common approach is totolerate the solubility of Li2Sn, at least in the positive electrodeenvironment, and protect the anode by, for example, interlayers,12,13 ion-selective separators,14,15 positive electrodes which
may actively adsorb polysuldes,16–18 or additives that aid in theformation of a favourable solid–electrolyte interphase (SEI).19–22
One of the most well-known redox shuttle-suppressing
strategies, lithium nitrate, LiNO3, generally used as an electrolyte co-salt or additive, falls into this latter category. It is
generally understood that the reduction of nitrate at the anodeand a subsequent oxidation of polysuldes aids in the formation of a denser SEI which helps to prevent the reduction ofpolysulde at the anode surface.20,22,23
While the use of LiNO3 and many other strategies have beenshown to considerably restrain the shuttle eﬀect and improvecycling, and approaches such as polysulde-impermeable layers
or solvents with low polysulde solubility have been validatedby measurements of sulfur content in the electrolyte, it is relatively rare that the electrochemical eﬃciency of cells is assessedby methods other than measurements of coulombic eﬃciency
from galvanostatic cycling, or other more qualitative measurements. However, a few detailed studies of redox shuttle and selfdischarge behaviour have been published.

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Mikhaylik and Akridge24 were the rst to quantitatively
describe the redox shuttle and its relation to capacity andcharge eﬃciency, through a combination of mathematicalmodelling and experiments covering the eﬀects of charging
rate, self-heating and salt concentration. In that work they
developed the concept of the “charge-shuttle factor”, a relationship between the eﬀective parasitic current to the appliedcurrent, and a shuttle rate constant to quantify the rate of self
discharge under rest conditions. More recently, Moy et al.25
quantied the redox shuttle by measuring the steady state
current passed through Li–S cells when held under potentiostatic control at pre-determined states-of-charge. The authors
demonstrated that the “shuttle current” is highest close to the
maximum state-of-charge, where the concentration of thehigher order (i.e., longer chain) polysuldes is highest. The
shuttle currents for cells containing LiNO3 in the electrolyte
were also measured, but were more than an order of magnitude
lower compared with the cells without LiNO3, and considered to
be almost negligible. Gordin et al.26 quantied self-discharge ofcells with diﬀerent electrolytes by inserting a 2 week rest period
between galvanostatic cycles at low rate, measuring thepercentage loss of capacity aer the rest period, while Hart
et al.16 similarly quantied self-discharge for a range of diﬀerent
sulfur host materials by stopping cells mid-cycle for three days
and measuring the capacity loss. Very recently, Xu et al.27
quantitatively assessed the eﬀect of self-discharge throughusing LiNO3 and a Naon-coated separator.
While LiNO3 (among other strategies) enables cycling of cells
at low rate with a reasonable coulombic eﬃciency, selfdischarge over even moderate storage or relaxation periods isstill a signicant issue and one that is largely unexplored. In
fact, as the development of the positive electrode continues toimprove and performance is increasingly benchmarked by ratecapability, the issue of self-discharge can easily be ignored andclaims of control over polysulde mass transport may not beproperly scrutinised.
In this work, we present a detailed overview of the redox
shuttle and self-discharge behaviour in electrolytes containing
LiNO3, through semi-quantitative measurements of polysuldes in the electrolyte when cells are at idle conditions,changes in the SEI, and electrochemical methods for quantifying self discharge over extended cycling. The aim of this work
is to demonstrate the signicant eﬀect of self-discharge on a Li–
S cell even when an additive considered eﬀective – as LiNO3 is –is used, even without particularly heavy usage of the cell. An
improved understanding of the redox shuttle and self-discharge
processes as well as the use of appropriate measurements toassess medium- and long-term charge storage is essential to thedevelopment of this system for wider practical application.

2 Experimental
2.1

Materials

Carboxymethylcellulose sodium salt (CMC, Leclanché), styrene
butadiene rubber (SBR, Targray PSBR-100), 1,2-dimethoxyethane (DME, Novolyte), 1,3-dioxolane (DOL, anhydrous,
Aldrich) and sulfur (S8, Aldrich) were used as received. LithiumThis journal is © The Royal Society of Chemistry 2016

RSC Advancesbis(triuoromethanesulfonyl)imide (LiTFSI, Novolyte), lithium
perchlorate (LiClO4, Aldrich), lithium nitrate (LiNO3, Aldrich)and lithium sulde (Li2S, Aldrich) were dried at 120  C undervacuum prior to use. A sample of a high conductivity carbon
black with a surface area of 1100 m2 g1 and a pore volume of
1.74 cm3 g1 as analysed and used in our earlier studies11,28 was
kindly provided by Orion Engineered Carbons GmbH.
2.2

Electrode preparation

Positive electrodes were prepared by rst mixing sulfur andcarbon black in a 58 : 35 ratio and heating to 155  C in order tomelt the sulfur into the pores of the carbon. The resulting
composite was then dispersed into water with the binders togive a composition of 58 : 35 : 7 S : C : binder, where the binder
was a 2 : 3 mixture of CMC : SBR. The slurry was mixed byplanetary ball-milling for 2 hours and coated onto Al foil toa loading of approximately 1 mgs cm2. The electrodes wereallowed to dry at ambient conditions, cut into circular discs ofthe desired size (13 mm or 20 mm in diameter), then transferred
into an Ar-lled glove box and dried further at 55  C overnight.
2.3

Cell preparation

Two diﬀerent cell preparation techniques were employed in thiswork. For all standard two-electrode electrochemical measurements on Li–S cells, CR2025 coin cells were assembled
comprising a 13 mm diameter cathode, a 17 mm diameterporous polyethylene separator (SOLUPOR, Lydall Performance
Materials) and a 16 mm diameter piece of a 125 mm-thick Li foil
(Cyprus Foote Mineral) as the anode. The electrolyte for all coincells was 1 M LiTFSI, 0.25 M LiNO3 in 1 : 1 DME : DOL, and theelectrolyte was xed at 6 mL mgs1.
Four-electrode pouch cells, including a second insulated
wire electrode as an in situ probe with common counter andreference electrodes, were prepared as described in our previous
work11 with some modications. Briey, the cells were constructed with a 20 mm diameter cathode, a 25 mm diameter
polyethylene separator on the positive electrode side and a 25
mm diameter glass bre separator (Whatman GF-A) on thenegative electrode side, and a 22 mm diameter piece of 125 mmthick Li foil as the negative electrode. The probe electrode wasa 125 mm diameter PTFE-insulated Pt wire (Advent ResearchMaterials). The electrolyte was the same as previously describedfor coin cell experiments, although a larger amount of electrolyte was required (35 mL mgs1) in order to achieve comparable
performance.
For analysis of anode surfaces, symmetrical LikLi pouch cells
were prepared using a 20 mm diameter piece of 125 mm-thick Li
foil as the working electrode to be analysed, a 22 mm diameter
polyethylene separator and a larger piece of Li foil as the anode.
The electrolyte was 80 mL of 1 M LiClO4, 0.25 M LiNO3 in 1 : 1
DME : DOL, presaturated with polysuldes in order to mimic
the electrolyte chemistry of the Li–S cell in the symmetricalsetup. Sulfur pre-saturation was achieved by agitating equimolar amounts of S8 and Li2S (giving a nominal stoichiometryof “Li2S9”) in the electrolyte for at least 3 days and ltering oﬀ
any remaining solid. LiClO4 was chosen as the electrolyte salt to

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avoid the inclusion of F-containing compounds in the SEI,
compounds which are highly susceptible to decomposition
under X-ray radiation. The use of LiClO4 also ensures that sulfur
compounds incorporated into the SEI can only originate fromthe dissolved polysuldes.

Paper

analysis. Samples were removed from the glove box in sealed
vials and then introduced to the microscope (Carl Zeiss SIGMA
FE-SEM) via a N2-lled glove bag.

3 Results and discussion

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2.4

Electrochemical testingElectrochemical measurements were made with either MPG2 or
VMP2 potentiostats (Bio-Logic). All galvanostatic cycling steps
were at a rate of C/10 (167.2 mA g s1) between the voltage limitsof 1.8 V and 2.6 V vs. Li/Li+. Three diﬀerent test protocols aredescribed in this work:(1) Shuttle current measurement under potentiostatic
control: Cells were cycled galvanostatically as described above.
Measurements were made on pre-determined cycles (1st, 5th,
10th and every following 10th cycle) by holding the cell potential
at 2.38 V on charge for 72 hours.
(2) Shuttle current measurement under OCV conditions(“cycle/wait” test): Cells were cycled galvanostatically as
described above. Every third cycle the cell was stopped at 100%charge (2.6 V) and le to relax at OCV for a period between 12hours and 14 days. Following the rest, cycling was recommenced with the discharge step.
(3) Polysulde concentration measurements with an in situ
probe: Tests were run on the VMP in bipotentiostat mode using
two synchronised channels to allow the use of common counterand reference electrodes. The cell was subjected to a single
discharge/charge cycle and le to relax under open circuit
conditions from 100% charged. Linear sweep voltammograms(LSVs) were then made at the probe electrode between OCV and2.85 V vs. Li/Li+ at a scan rate of 1 mV s1 every four hours for
200 hours.
For symmetrical LikLi cells, “discharge” and “charge”
currents (400 mA cm2) were calculated based on the area of the
20 mm diameter Li working electrode. The discharge andcharge times were 6 hours each for all experiments (i.e., 2.4 mA
h cm2).
2.5

Surface analysis

Cells were disassembled in an Ar-lled glove box. All extracted
samples were washed with a 1 : 1 mixture of DME : DOL.
Samples analysed by X-ray photoelectron spectroscopy (XPS)
were rst transferred to the spectrometer through a sealedtransfer system to avoid air exposure. XPS characterisation was
performed using a PHI 5500 spectrometer (Physical Electronics)
using monochromatic Al K radiation (1468.7 eV). Spectra wererecorded with a pass energy of 23.5 eV giving an overall
instrumental resolution of 0.6 eV as determined from thebroadening of a Ag Fermi edge. All sulfur lines are curve tted
using a doublet peak with a spin–orbit splitting of 1.18 eV andan intensity ratio of 2 : 1 characteristic for S2p. Spectra were
calibrated in binding energy to the hydrocarbon peak at 285 eV.
Reference samples of Li2SO4, Li2S, Na2S2O3 and Na2S2O3 were
measured to assist with peak assignment (ESI, Fig. S1†).
Following XPS measurements the same transfer system wasused to return the samples to the glove box to prepare for SEM

3634 | RSC Adv., 2016, 6, 3632–3641

3.1“Baseline” Li–S cell electrochemistry

For studying the behaviour of “full” Li–S cells, we have useda simple cathode composition similar to that used in ourprevious work11,28 comprising a porous carbon/sulfur composite
with a total sulfur content of 58% in the electrode. The electrolyte volume used, at 6 mL mgs1 was rst determined as thelowest possible amount which could be used with the cathodes
without resulting in a serious decrease in cell performance. For
comparison purposes, the galvanostatic cycling performance ofthe benchmark cell used in this work is given in Fig. 1.
Aer an initial drop in capacity over the rst few cycles, thecell gives a relatively stable reversible capacity of 800 mA h g1
over the rst 100 cycles. The coulombic eﬃciency, however, isrelatively low, decreasing from 95% in the initial cycles to92% over the rst 100 cycles. The low coulombic eﬃciency is
expected given that no specic strategies for suppressing theredox shuttle other than the addition of LiNO3 to the electrolyte
are used here, and the use of the smallest possible amount ofelectrolyte ensures that the electrolyte becomes saturated with
polysuldes during the charge and discharge process, ensuring
the highest possible rate of reaction with the anode surface.
Furthermore, the sulfur loading in the cell is relatively low at
1 mgs cm2, which as previously reported26 may in fact
contribute further to the rate of self-discharge. The continuous
decrease in the coulombic eﬃciency can be interpreted asa continuous increase in the rate of self-discharge over thecycling period. However, as has been previously discussed,25 thecoulombic eﬃciency does not give any specic information
about the rate of self-discharge and can vary considerably
depending on sulfur utilisation, charging rate and so on.
Nonetheless, this behaviour is typical of cells of this type
without a specically optimised positive electrode and isa reasonable baseline for further investigation.

Fig. 1 (Bottom) Discharge capacity vs. cycle number for a Li–S cell
cycled galvanostatically at a constant rate of C/10 (167.2 mA gs1); (top)coulombic eﬃciency (Qdc/Qc) vs. cycle number for the same cell.

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3.2

RSC Advances

Electrolyte–anode interface and SEI formation

The reactions of dissolved polysuldes with Li metal and thesurface layer formed on the anode in contact with the electrolyte(without cycling, with cycling, and aer a long relaxation time
following cycling) was studied with X-ray photoelectron spectroscopy (XPS) using simplied symmetrical (LikLi) cells with anelectrolyte pre-saturated with polysuldes and a supporting salt
not containing S (LiClO4). SEM images for Li metal samples are
given in Fig. 2, and XPS spectra for the S2p orbital are given in
Fig. 3.

Fig. 3 S2p XPS spectra for Li metal anode samples. (a) Without
exposure to the electrolyte (pristine Li); (b) Li soaked in the electrolyte
for 12 h (soaked Li); (c) Li after 1 discharge/charge cycle in the polysulﬁde-saturated electrolyte (charged Li); (d) Li cycled as in (c) and left
to rest for 10 days before extraction from the cell (self-disch.).

SEM images of Li anodes. (Top) Pristine Li, (middle) after 1
discharge/charge cycle (“charged”), (bottom) after 1 cycle and 10 days
rest (“self-discharged”).

Fig. 2This journal is © The Royal Society of Chemistry 2016

As is clear from Fig. 3a, the as-received Li foil is free from any
S-containing compounds in the surface layer. Following 12hours of soaking in the polysulde-saturated electrolyte, several
diﬀerent S-containing species are present (Fig. 3b). The mostprominent emission at 169.2 eV coincides in binding energy
with sulfate ([SO4]2, S(+VI)). The peak at 167.2 eV originatesfrom the S(+IV) oxidation state, and likely assignments are
sulte ([SO3]2) or the central S-atom in lithium thiosulfate([S–SO3]2). These oxidised sulfur species are an expected resultof oxidation reactions following reduction of the LiNO3 additivein the electrolyte, as has been investigated previously.20 Thiosulfate is a likely product of a disproportionation reaction
between sulte and polysulde. For thiosulfate, a peak for[S–SO3]2 – should be accompanied by a second peak at 162 eV
for [S–SO3]2 in a 1 : 1 ratio29 (see also the reference spectra in
the ESI, Fig. S1†).
In this case, the peak observed in this binding energy range
is smaller in intensity. Terminal atoms (ST) of polysuldespecies (Li2Sn, e.g., Li2S2 and Li2S4) may furthermore also
contribute to the photoelectron intensity in this binding energy
range. Therefore, it is most likely that the peak at 167.2 eV isprimarily from [SO3]2. The remaining peak at 164 eV is

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attributed to the S(0) oxidation state. Since this peak is higher inintensity than the peak at 162 eV (ST), it is plausible that thissignal originates from the bridging-S (SB) of metastable Li2Sn
where n > 4. However, just as it is diﬃcult to distinguish[S–SO3]2 from ST in Li2Sn, it is diﬃcult to distinguish an SB
environment from elemental sulfur (S8). It therefore cannot be
currently established if higher order Li2Sn exists in the surface
layer or if such compounds have decomposed into S8 and lower
order Li2Sn; though not, as evidenced by a lack of a peak at
160 eV, lithium sulde (Li2S).
Aer a single discharge/charge cycle with immediate
extraction from the cell, SEM analysis reveals the expected“mossy” deposition of lithium metal (Fig. 2, middle), and ve
diﬀerent sulfur environments can be distinguished in the XPS
spectrum (Fig. 3c). Of principal importance is the appearance ofan emission from Li2S (S(II)) at 160 eV. Relative to the“soaked” sample (Fig. 2b), the “charged” sample also showsincreased intensity in the peaks at 167 eV ([S–SO3]2 or [SO3]2)and 162 eV ([S–SO3]2 or ST).
Aer 10 days exposure to the electrolyte following a single
discharge/charge cycle (Fig. 3d), the S2p spectrum is dominatedby the emission located at 162 eV (ST and [S–SO3]2). Comparedto the “soaked” and “charged” samples, the emission of SB isshied to lower binding energy by 0.7–0.9 eV. This may indicate the presence of an environment such as the SB-atom in S32(i.e., [S–S–S]2), in which SB is more strongly aﬀected bysecondary chemical shis as compared with longer polysuldes. Crucially, contributions from Li2S and [SO4]2 arealmost negligible, indicating a general increase in thickness of
the SEI which obscures these compounds, since the probing
depth of XPS is only of the order of a few nanometres. Thisobservation is supported further by the appearance of insulating particles on the surface as seen with the SEM (Fig. 2,
bottom). Li2S, perhaps surprisingly, appears to be added to thesurface layer only in very small quantities while the cell is not in
operation. Given the overall intensities of the peaks at 162 eV
and 167 eV it is most likely that the major products formed on
the anode during a long idle time are short chain polysuldessuch as Li2S2 and possibly Li2S3, with additional contributions
from oxidised sulfur species such as Li2SO3 and Li2S2O3 formedby reaction of the reduction products of LiNO3.
It should be noted here that the substitution of LiTFSI for
LiClO4 could be expected to aﬀect the reactions of polysuldesat the negative electrode surface; LiClO4, like LiNO3, isa powerful oxidising agent which may contribute to thepassivation of the negative electrode. Indeed, a shuttleinhibiting eﬀect from the use of LiClO4 has been previously
reported by Kim et al.30 However, in that work, while the use ofan electrolyte of 1 M LiClO4 oﬀered some improvement incoulombic eﬃciency compared to 1 M LiTFSI (80% comparedto 70%), an electrolyte containing 0.2 M LiNO3 in addition to
LiClO4 showed by far the highest coulombic eﬃciencies (95–100%; 1 M LiTFSI and 0.2 M LiNO3 was not tested, but our own
tests under similar conditions gave similar results). LiNO3 ismuch less kinetically stable than LiClO4, and this is the most
likely reason for its stronger eﬀect in passivating the negativeelectrode. Consequently, we believe the diﬀerences in the3636 | RSC Adv., 2016, 6, 3632–3641

Paperreactions of polysuldes at the negative electrode in the presence of LiTFSI or LiClO4 to be minor while a large amount of
LiNO3 is present, and that the conclusions we draw from theseexperiments remain valid for the LiTFSI electrolyte system.

3.3 Changes in electrolyte composition during selfdischarge
The processes governing self-discharge can be visualised
through the use of an in situ electrochemical probe, as we havepreviously reported.11 The technique uses a small insulated wire
(probe) electrode placed between two separators in an otherwise
standard cell setup. The probe electrode and cell positive electrode can be simultaneously controlled or measured using
common counter and reference electrodes. At the probe electrode, standard electrochemical techniques can be used toanalyse the behaviour of polysuldes dissolved in the electrode,outside of the positive electrode environment.
In this work, we used a PTFE-insulated Pt wire cut to expose
a 125 mm-diameter Pt surface inside the cell. To allow forincorporation of the probe electrode, a pouch cell format rather
than a coin cell was used. For this reason, and because of theneed for an extra separator, the cell requires a considerablylarger volume of electrolyte (35 mL mgs1 was found to give
approximately the same cell performance when comparing
pouch cells with the coin cell baselines). Despite these diﬀerences, the reactions occurring in solution are not expected to besignicantly diﬀerent and we consider observations made withthis technique to be valid at the very least for qualitative
purposes.
To assess the changes in polysulde concentration in theelectrolyte, linear sweep voltammetry (LSV) measurements weremade at 4 hour intervals over the course of a long (>200 hour)relaxation period following a single discharge/charge cycle. The
results have been summarised in Fig. 4. Each LSV measurementshows either a peak or a plateau in current at around 2.75 V vs.
Li/Li+, consistent with the oxidation of all polysulde species toelemental sulfur. Since the peak or steady-state current ina diﬀusion-limited voltammetric measurement at a planarelectrode is proportional to the concentration of the electrochemically active species, we can take the peak current in thiscase as an indicator of the concentration. However, it is notpossible to ascribe it directly to relative concentrations of polysuldes, because the number of electrons transferred perpolysulde is not constant; that is, if the major polysuldespecies in solution are S42 and S62, then the major electrochemical oxidations to elemental sulfur would be:
2S4 2 /S8 þ 4e
4 2
8
S6 /S8 þ e
3
3

(1)

Since, as is known, the ratio of S62 and S42 (among otherspecies) varies during the charge/discharge process, thenumber of electrons therefore also varies. Nonetheless,consideration of the peak current is still qualitatively useful.

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Results of measurements conducted with a 125 mm diameter Pt in situ probe electrode over a 200 hour OCV relaxation period of a Li–S
cell from 100% charged. (Left) Linear sweep voltammograms (LSVs) for the probe electrode at a scan rate of 1 mV s1. (Right-top) Cell OCV vs.
relaxation time, (right-bottom) peak currents (jmax) from LSVs at the probe electrode vs. relaxation time.
Fig. 4The peak currents from the LSVs in (Fig. 4 le) were extracted
and plotted against time alongside the cell OCV (Fig. 4 right).
From the charged state, the peak current increases gradually
over the rst 125 hours indicating a gradual increase in theconcentration of polysuldes in the electrolyte. The largest
value for the peak current coincides almost exactly with theminimum in the cell OCV (i.e., the clear start of the lower
voltage plateau). This is consistent with what is already known
about the discharge process; the lower voltage plateau even at
low discharge rate generally begins with a peak in the voltage,
usually ascribed to an increase in overpotential caused by oversaturation of polysuldes.
The upper plateau is understood in this electrolyte system tocorrespond to the complete conversion reduction of S8 with S62as the major product:
84S8 þ e / S6 2
33

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(4)

Combination of eqn (3) and (4) gives an overall reaction of:
4Li + S8 ! 2Li2S4(5)

in which S62 acts as the mediating or shuttling species.
Aer 125 hours, the cell voltage remains roughly constant at
close to 2.1 V, and the peak currents as measured by the probe
begin to decrease, indicating a decreasing concentration ofsoluble polysuldes. This could indicate further consumptionof polysuldes by the anode, forming further reduced species
such as Li2S2 or Li2S3, as indicated previously by XPS. Alternatively, it may be that lower-order polysuldes such as Li2S4 tends
either to precipitate as a solid or disproportionate into othersolids such as Li2S and S8 elsewhere in the cell.

(2)

The coincidence of the start of the lower plateau – implying
an equilibrium of S62 and S42 in the electrolyte – and the
maximum in the peak current suggests that at this point allelemental sulfur formed on the charge has been re-reduced
back to soluble forms. This process is driven by the redox
shuttle eﬀect as it is generally understood, where lower orderpolysuldes formed by reaction with the anode can react withelemental sulfur – and not necessarily be oxidised only byelectron transfer. Such reactions could be written as, forexample:2Li + 2S62 ! 2Li+ + 3S42

4S42 + S8 ! 4S62

(3)

3.4 Quantifying self-discharge and its eﬀect on cycling
performance
Self-discharge can be simply quantied by measuring thecapacity of the cell before and aer a specied rest time. In thiswork, we have measured self-discharge for multiple relaxation
periods ranging from 12 hours to 2 weeks in duration (hereaer
referred to as a “cycle/wait” test for simplicity). Studying a rangeof relaxation times over an extended period enables thebuilding of a detailed picture of cell stability when the cell isboth in and not in continuous operation. The results of thisexperiment are summarised in Fig. 5.
As seen in Fig. 5a, the cell shows an initial capacity in excess
of 1000 mA h g1 which decreases over the rst 10 cycles toa relatively stable reversible capacity of about 800 mA h g1,

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Fig. 5 Self-discharge analysis by a “cycle/wait” test: (a) capacity vs.
cycle number, where every dark red point is the discharge capacity
following a wait of the indicated number of days since the previous
cycle. (b) Calculated capacity loss values from (a), over the total rest
period (in black) and the estimated relative capacity loss (i.e., thederivative of capacity loss wrt relaxation time, in grey). Cycling rate is
C/10 (167.2 mA g1).

which is consistent with the reference cell (Fig. 1). The charge
eﬃciency also remains at >90% throughout the experiment also
consistent with the reference cell. It is also easily seen from thisplot that the capacity loss increases with increasing OCV
relaxation time and that the capacity loss does not increasesignicantly aer the rst three days. It is therefore immediately apparent that the rate of self-discharge changes dependingon the state of charge, as has been previously noted by Moy
et al.25 The capacity loss for each rest period was calculated as
the diﬀerence between the measured discharge capacity and theestimated discharge capacity in the absence of a rest period(extrapolated from the discharge capacity of the cycles immediately before and aer). This analysis is shown graphically in
Fig. 5b. An alternative plot presenting the cell voltage proles on
the same time axis as the capacity loss is given in the ESI,
Fig. S2.† In order to put the rate of self-discharge into context,
one can dene an “equivalent C-rate”, i.e., the current thatwould have to pass in an ideally eﬃcient cell to draw the same
capacity in the same time. This can be simply calculated fromthe capacity loss divided by the relaxation time, expressed asa fraction of the theoretical capacity:

expressed as equivalent C-rates over both the whole relaxation
times (total rest) and between relaxations of diﬀerent times(relative). In the relative case, for example, the diﬀerence incapacity losses from rest times of 24 h and 72 h are compared inorder to estimate the equivalent rate over the last 48 h of the 72
h rest. These values are given in Table 1.
From Fig. 5b and Table 1 it is easily seen that self-discharge
is fastest during the rst three or so days and subsequently
slows dramatically. Furthermore, the rate of self-discharge
appears to increase as the cell ages, as indicated by the larger
capacity losses on cycles 17–32 compared with the rst 16
cycles. Over the rst 24 hours, the rate of self-discharge is of theorder of C/300, approximately 30 times slower than the charging
rate. Aer the rst 16 cycles this rate increases to a value of theorder of C/200. Aer the rst 72 hours the rate of discharge
sharply decreases to a point at which it is too slow to bemeasured accurately (i.e., changes of only a few mA h g1 overthe course of several days). A maximum capacity loss of approx.
278 mA h g1 could be considered consistent with the interpretation that the capacity loss comes largely from the reduction of elemental sulfur in the cathode, since the rst reductionof sulfur:
84
S8 þ e / S6 2
3
3

(7)

has a theoretical capacity of 278 mA h g1. However, since it has
previously been shown that some elemental sulfur may remain
inactive in the electrode during cycling,31 it is likely that thiscapacity loss contains contributions from the reduction ofsoluble polysuldes (e.g., S62 to S42).
For further comparison, the capacity losses for relaxation
times up to three days were analysed according to the model
proposed by Mikhaylik and Akridge,24 in which a shuttle
constant can be derived according to the equation:ln

QH¼ ks trelax
Q0H

(8)

QH
is the remaining fraction of the upper plateau (Q0H in
Q0H
this case was taken to be 278 mA h g1), trelax is the relaxation

where

Table 1 Analysis of redox shuttle/self-discharge rate from the “cycle/
wait” test expressed as “equivalent C-rate”, over the total rest times orrelative to the previous rest

Rest time

C-Rate equiv.
(total rest)

C-Rate equiv.
(relative)

(6)

12 h
24 h
72 h
7d
14 d

C/330
C/280
C/510
C/1200C/2000C/330 (12 h)
C/240 (12 h)
C/890 (48 h)
C/29 000 (4 d)C/8400 (7 d)where C ¼ 1672 mA h g1. As has already been discussed, therate of self-discharge is of course not constant, but it is still
useful to dene the rate in normalised terms over specicrelaxation times. The capacity losses from Fig. 5b have been

12 h
24 h
72 h
7d
14 d

C/175
C/210
C/440
C/990
C/2000C/175 (12 h)
C/260 (12 h)
C/960 (48 h)
C/18 000 (4 d)—equiv: C-rate; C=n ¼

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time and ks is the shuttle constant, with units of h1. The
shuttle constant can then be determined by the gradient of the
QHline in a plot of ln 0 vs. relaxation time, which is given in the
QH
ESI, Fig. S3.† From this analysis we arrive at shuttle constants of
0.027 h1 for the rst loop in this experiment (cycles 1–16), and
0.049 h1 for the second loop (cycles 17–32). These values are
somewhat higher than the values of 0.014 h1 as very recently
reported for similar electrolyte systems.27
That self-discharge appears to eﬀectively cease aer the cells
reach the lower voltage plateau, coupled with the observationfrom XPS that Li2S is not formed to any signicant extent on thenegative electrode during self-discharge, may suggest that there
is an equilibrium of soluble polysuldes with the anode, as ineqn (3), i.e.:
2Li + 2S62 # 2Li+ + 3S42

(9)measurements made at the 1st, 5th and 10th cycles are given in
Fig. 6.
The current transients show a decay in the current toapproximately 0.7 A cm2 (corresponding approximately to an
C-rate equivalent of C/2500) following the rst cycle. This steadystate current increases to approximately 3 A cm2 on the hcycle, and gradually decreases with continued cycling (transients up to the 60th cycle are given in the ESI, Fig. S4†). Since
the rate of the redox shuttle would be expected to be proportional to both the surface area of the anode and the concentration of polysuldes at the surface, it is most likely that theincrease in the redox shuttle current from the rst to the hcycle is the consequence of the increasing surface area of thelithium anode over the rst few cycles. The increasing roughness of lithium within a small number of cycles can be seenquite clearly with SEM, as shown in Fig. 7. The decreasing
shuttle current over continued cycling could then be a result

and that the reduction of polysuldes further to compounds
such as Li2S is kinetically very slow.
An important observation is that this self-discharge is verymuch reversible: even aer two weeks and two months oftesting the cell is still capable of delivering the same capacity of
800 mA h g1 as the reference cell. This good reversibility may
be a direct result of the slow or even absent conversion of polysuldes to more kinetically inert compounds such as Li2S.

3.5

Redox shuttle measurement under potentiostatic control

The rate of the redox shuttle itself, as opposed to the rate ofcapacity loss, can be assessed by measurement of the current
passing through a cell under potentiostatic control. Thismethod was previously used by Moy et al.25 to determine thevariation in the eﬀective redox shuttle current with the cellstate-of-charge. Here, we have investigated the change in theredox shuttle on cycling by holding the cell potential at thearbitrary value of 2.38 V for three days on selected charges. The

SEM images of Li metal electrodes from symmetrical LikLi cells
subjected to a number of 6 hour discharge/6 hour charge cycles ata current density of 400 mA cm2.

Fig. 7

Fig. 6 Cell voltage and current vs. time for a Li–S cell subjected to a “shuttle current” measurement. On the 1st, 5th, 10th and every following
10th cycle the cell is charged to 2.38 V and held at this potential for 72 h.

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either of a decreasing concentration of polysuldes, if, forexample, decomposition of the electrolyte reduces the solubility, or if an increasingly thick SEI reduces the rate of thereaction between lithium and polysuldes.
It should be noted that the reduction of the shuttle current
over cycling as determined by this method conicts with the“cycle/wait” measurements discussed previously, which showed
an increased rate of self-discharge with increased cycling. Whilethe diﬀerences (i.e., increases or decreases in the rate of selfdischarge) are relatively minor, this nonetheless indicates that
self-discharge and lithium anode stability could be signicantlyinuenced by real-life usage, for example charging rate, storage
conditions, etc. To the best of our knowledge, such a study has
not yet been undertaken, and would be an interesting direction
for future work.

4 Conclusions
In this work, we have explored the self-discharge behaviour ofsimple Li–S cells containing the common electrolyte additive
LiNO3, and related this behaviour to the polysulde redox
shuttle phenomenon through a range of electrochemical andsurface analysis techniques. LiNO3 is an eﬀective suppressant of
the polysulde redox shuttle – reducing the eﬀective electronic
leak by one or two orders of magnitude – and enables thereversible cycling of Li–S cells with coulombic eﬃciencies inexcess of 90% at low cycling rates. However, the redox shuttle isstill fast enough to cause cells to self-discharge at a relatively
rapid rate, with cells in this work observed to lose more than
25% of their capacity from fully charged in less than three days.
This capacity loss is shown to largely come from the reductionof elemental sulfur back to soluble polysuldes, as observed
through the use of an in situ electrochemical probe. Thechanges in anode surface chemistry aer self-discharge wasinvestigated with XPS, where a signicant increase in theamount of reduced sulfur species was found aer long exposure
to the electrolyte following cycling.
A key observation is the relatively good reversibility of theself-discharge of these cells; that is, cells rested for weeks
following a full charge were still able to deliver the samedischarge capacities as those which were cycled continuously.
This reversibility is attributed primarily to the relative lack ofcontinuous formation of Li2S on the anode during selfdischarge, as conrmed by XPS, which in turn is due to thevery slow reaction of polysuldes with the anode once the cellvoltage has reached equilibrium at the lower voltage plateau of2.15 V.
From a practical perspective, the rapid self-discharge in thissystem, despite the application of an electrolyte considered tobe eﬀective, remains a challenge which merits considerably
more attention, especially if the system is expected to challenge
existing Li-ion batteries in any consumer application. Furthermore, it is important for future research into redox shuttle
suppression strategies that self-discharge be quantied directly,
by methods such as, for example, the “cycle/wait”-type
measurements as discussed in this work, rather than assessedindirectly from the coulombic eﬃciency on galvanostatic

3640 | RSC Adv., 2016, 6, 3632–3641

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cycling. This is especially important if cells are tested at relatively fast rates, since self-discharge is increasingly uninuential the faster cells are continuously cycled.

Acknowledgements
The SEM and XPS analysis included in this work in Sections 3.2
and 3.5 formed part of a Master degree project entitled “Insights
into the morphological changes undergone by the anode in thelithium sulphur battery system” by Anurag Yalamanchili under thesupervision of Dr Matthew Lacey and Dr Carl Tengstedt in 2014. The
full text of the nal thesis is available online from Uppsala University
at http://www.diva-portal.org/smash/record.jsf?pid¼diva2:765276.
The authors wish to thank Dr Martin Cadek (Orion Engineered
Carbons GmbH) for providing the carbon black used in thiswork, and the Era Net Transport project “MaLiSu” and Vinnova
in Sweden for nancial support.

References
1 A. Manthiram, S.-H. Chung and C. Zu, Adv. Mater., 2015, 27,
1980–2006.
2 A. Manthiram, Y. Fu and Y.-S. Su, Acc. Chem. Res., 2013, 46,
1125–1134.
3 D. Bresser, S. Passerini and B. Scrosati, Chem. Commun.,2013, 49, 10545.
4 M. A. Pope and I. A. Aksay, Adv. Energy Mater., 2015, 1500124.
5 K. Ueno, J. W. Park, A. Yamazaki, T. Mandai, N. Tachikawa,
K. Dokko and M. Watanabe, J. Phys. Chem. C, 2013, 117,20509–20516.
6 M. Cuisinier, P.-E. Cabelguen, B. D. Adams, A. Garsuch,
M. Balasubramanian and L. F. Nazar, Energy Environ. Sci.,
2014, 7, 2697.
7 J. T. Lee, Y. Zhao, S. Thieme, H. Kim, M. Oschatz,
L. Borchardt, A. Magasinski, W.-I. Cho, S. Kaskel andG. Yushin, Adv. Mater., 2013, 25, 4573–4579.
8 E. S. Shin, K. Kim, S. H. Oh and W. Il Cho, Chem. Commun.,2013, 49, 2004–2006.
9 L. Suo, Y.-S. Hu, H. Li, M. Armand and L. Chen, Nat.
Commun., 2013, 4, 1481.
10 Y.-C. Lu, Q. He and H. A. Gasteiger, J. Phys. Chem. C, 2014,
118, 5733–5741.
11 M. J. Lacey, K. Edström and D. Brandell, Electrochem.
Commun., 2014, 46, 91–93.
12 X. Wang, Z. Wang and L. Chen, J. Power Sources, 2013, 242,
65–69.
13 Y. S. Su and A. Manthiram, Nat. Commun., 2012, 3, 1166.
14 I. Bauer, S. Thieme, J. Brückner, H. Althues and S. Kaskel, J.
Power Sources, 2014, 251, 417–422.
15 Z. Jin, K. Xie, X. Hong, Z. Hu and X. Liu, J. Power Sources,
2012, 218, 163–167.
16 C. J. Hart, M. Cuisinier, X. Liang, D. Kundu, A. Garsuch and
L. F. Nazar, Chem. Commun., 2015, 51, 2308–2311.
17 Q. Pang, D. Kundu, M. Cuisinier and L. F. Nazar, Nat.
Commun., 2014, 5, 1–19.

This journal is © The Royal Society of Chemistry 2016

View Article Online

Published on 15 December 2015. Downloaded by Uppsala University on 18/07/2017 13:47:16.

Paper

18 J. Song, T. Xu, M. L. Gordin, P. Zhu, D. Lv, Y.-B. Jiang,
Y. Chen, Y. Duan and D. Wang, Adv. Funct. Mater., 2013,
24, 1243–1250.
19 C. Zu and A. Manthiram, J. Phys. Chem. Lett., 2014, 5, 2522–
2527.
20 D. Aurbach, E. Pollak, R. Elazari, G. Salitra, C. S. Kelley and
J. Aﬃnito, J. Electrochem. Soc., 2009, 156, A694.
21 Z. Lin, Z. Liu, W. Fu, N. J. Dudney and C. Liang, Adv. Funct.
Mater., 2013, 23, 1064–1069.
22 S. S. Zhang, Electrochim. Acta, 2012, 70, 344–348.
23 W. Li, H. Yao, K. Yan, G. Zheng, Z. Liang, Y.-M. Chiang andY. Cui, Nat. Commun., 2015, 6, 7436.
24 Y. V. Mikhaylik and J. R. Akridge, J. Electrochem. Soc., 2004,
151, A1969–A1976.
25 D. Moy, A. Manivannan and S. R. Narayanan, J. Electrochem.
Soc., 2014, 162, A1–A7.

This journal is © The Royal Society of Chemistry 2016

RSC Advances

26 M. L. Gordin, F. Dai, S. Chen, T. Xu, J. Song, D. Tang,
N. Azimi, Z. Zhang and D. Wang, ACS Appl. Mater.
Interfaces, 2014, 6, 8006–8010.
27 W.-T. Xu, H.-J. Peng, J.-Q. Huang, C.-Z. Zhao, X.-B. Cheng
and Q. Zhang, ChemSusChem, 2015, 8, 2892–2901.
28 M. J. Lacey, F. Jeschull, K. Edström and D. Brandell, J. Phys.
Chem. C, 2014, 118, 25890–25898.
29 X. Liang, C. Hart, Q. Pang, A. Garsuch, T. Weiss and
L. F. Nazar, Nat. Commun., 2015, 6, 5682.
30 H. S. Kim, C.-S. Jeong and Y.-T. Kim, J. Appl. Electrochem.,
2011, 42, 75–79.
31 M. Cuisinier, P.-E. Cabelguen, S. Evers, G. He, M. Kolbeck,
A. Garsuch, T. Bolin, M. Balasubramanian and L. F. Nazar,
J. Phys. Chem. Lett., 2013, 4, 3227–3232.

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