A Brief Review: Past, Present and Future of Lithium Ion Batteries 1

Transcription

1 ISSN , Russian Journal of Electrochemistry, 2016, Vol. 52, No. 12, pp Pleiades Publishing, Ltd., Published in Russian in Elektrokhimiya, 2016, Vol. 52, No. 12, pp A Brief Review: Past, Present and Future of Lithium Ion Batteries 1 Florian Schipper* and Doron Aurbach Bar-Ilan University, Department of Chemistry, Ramat-Gan, Israel * florian.schipper@gmail.com Received March 28, 2016 Abstract The review summarizes the development of lithium ion batteries beginning with the research of the s which lead to modern intercalation type batteries. Following the history of lithium ion batteries, material developments are outlined with a look at cathode materials, electrolyte solutions and anode materials. Finally, with lithium sulfur and lithium oxygen batteries two post intercalation type lithium batteries are discussed. The focus of the material discussions lies on basic understanding, problems and opportunities related to the materials. Keywords: lithium ion batteries, intercalation electrodes, cathode materials, anode materials, lithium-sulfur batteries, lithium-oxygen batteries DOI: /S INTRODUCTION Electricity is a ubiquitous commodity and life as we know it is hard to imagine with an absence of convenient power supply. The first battery ever conceived was invented by Alessandro Volta in 1800 to study the findings of Luigi Galvani on animal electricity [1]. The so called voltaic pile was crucial to early electrochemical experiments paving the way to fundamental understandings in the field and the synthesis of pure alkali metals was achieved via electrolysis powered by voltaic piles [2]. The primary Leclanche cell, a predecessor to the zinc carbon cell of today, was used in telegraphy stations and early telephones making it one of the first commercial batteries [3]. With the invention of the electrical generator in the mid-19th century the need for a rechargeable storage system arose and was met by the lead acid battery of Gaston Plante [4]. Interestingly, combustion engines and electrical engines competed for some time as means of propulsion in cars around the turn of the 19 20th century [5, 6]. Unfortunately, with more available petroleum reserves discovered in the early 20th century and the development of more comfortable gasoline cars, electric vehicles were soon outperformed by internal combustion engines and discontinued. Nonetheless, lead acid batteries are still a crucial part in gasoline powered cars as starter battery, making them the most successful secondary battery of the last century. New impulses to the battery field came in the late 20th with the development of portable consumer electronics (cell phones, camcorders, portable computers) 1 The article is published in the original. which require a battery of high energy density. The lead acid battery was not suitable for such applications and much attention was given to the nickel cadmium system which was eventually overtaken by the lithium ion batteries (LIB) in the early 1990s. LIB, as a part of our daily life, have been improved constantly over the last 25 years. They have attracted tremendous research attention with more than papers mentioning LIB in 2015 alone. Followed by the huge successes of LIB other battery systems (e.g. magnesium and sodium) have seen increased research interested as well making the field ever more fruitful [7, 8]. LIB are connected to big promises made by industry and scientists to play a vital role in many fields ranging from the development of electrical vehicles, load levelling to energy storage of alternative power sources (solar, wind etc.). With some author even going as far as to foreshadowing resource wars hinging on the development of adequate batteries [9]. The text at hand is meant to be a short overview of the lithium battery field aimed at new students of the field or interested readers of neighboring disciplines. Emphasis is given to the development of LIB with a short historical breakdown of the exciting work of the 1970s 1980s. The second chapter deals with some of the more popular cathode intercalation compounds and will deal with basics, problems and opportunities of new intercalation materials. Electrolyte solutions and anode materials will be covered briefly at the end of the second chapter. The closing chapter is devoted to technologies beyond LIB and a very brief description of the lithium sulfur and lithium oxygen battery will be given. 1095

2 1096 FLORIAN SCHIPPER, DORON AURBACH (а) 3.0 (b) Li OR Cell emf, V ma 10 ma Structure of TiS 2 and LiTiS x in Li x TiS Fig. 1. (a) Intercalation of lithium into the layered structure of TiS 2, reprinted with permission from ref. [14], copyright Elsevier (b) Discharge voltage profile of Li x TiS 2 in a solution of lithium aluminum chloride in methyl chloroformate, reprinted with permission from ref. [15], copyright Science THE WAY TO LITHIUM ION BATTERIES ( ) Before the discovery of lithium intercalation into host structures mainly primary lithium batteries were of research interest. The investigation of high temperature molten secondary lithium batteries (e.g. Li/S and Li/Cl 2 ) never amounted to any commercial success due to cost and problems with the inherent highly corrosive system [10]. Research into the field of molten lithium systems was soon discontinued in favor of the more promising sodium sulfur system [11]. However, the combination of a low molecular weight (M Li :7 g mol 1 ), low density (ρ Li : g cm 3 ) with a very low standard electrode potential (E 0, Li : 3.04 V vs. standard hydrogen electrode) makes metallic lithium an attractive choice as an anode material in batteries. The low potential also means that electrolyte solutions will react with metallic lithium and early primary cells focused on solid electrolytes to overcome this issue. Unfortunately, solid electrolytes don t allow for high currents and all solid state primary lithium batteries were not successful in a wide range. The solid state lithium lithium iodine iodine (Li LiI I 2 )-system has been a success in cardiac pace makers since the early 1970 s and is still in use today allowing for a compact design of pace makers [12, 13]. In 1975 Wittingham et al. demonstrated via a wet chemical route (employing n-butyl lithium) the intercalation of lithium into a variety of layered transition metal dichalcogenides (sulfides and selenides) [14]. Lithium is intercalated between the transition metal layers at a one mole ratio leading to an expansion of the c lattice parameter (Fig. 1a). Out of the range of studied compounds TiS 2 was particularly interesting since it offers the lowest molecular weight, potentially lowest cost and is an electronic conductor. The electrochemical intercalation of lithium into TiS 2 demonstrates an average intercalation voltage of ca. 2 V vs. Li + /Li reaching nearly the theoretical capacity of 240 ma h g 1 (Fig. 1b) [15]. The idea of lithium intercalation batteries was picked up by EXXON and turned into an attempt at secondary lithium batteries in non-aqueous liquid electrolyte solutions [16]. One major drawback of TiS 2 is the rather elaborate synthesis. The material has to be kept under inert conditions at all time to avoid the decomposition to titanium oxide and hydrogen sulfide upon the reaction with water. Another problem arises from the metallic lithium anode itself. Early studies into non-aqueous liquid electrolyte solutions in combination with alkali metals showed the formation of a surface layer. The so called solid electrolyte interface (SEI) is conductive for alkali metals but is an electrical insulator [17]. In general the SEI formation is beneficial and will passivate the lithium surface making metallic lithium stable in water free organic electrolyte systems. Unfortunately, the SEI is not stable during prolonged cycling and will crack open which leads to a continuous consumption of electrolyte and lithium for the continuous SEI reformation [18]. This leads to rater low Coulomb efficiency for lithium of around 90% and secondary metallic lithium batteries have to be designed with an adequate excess of lithium metal [18]. Even worse, the uneven deposition of lithium onto the cracked SEI leads to the growth of lithium dendrites. The dendrites will eventually break off and form nano-sized islands of highly reactive lithium effectively lowering the thermal stability of the cell [19]. In order to overcome the issue of metallic lithium anodes attempts at lithium alloys were made, e.g. Li/Al [20], Li/Mg [21]. Using lithium alloys may lower the dendritic growth rate but at the cost of capacity and electrode potential. For Li/Al the deposition potential of lithium is around 0.4 V vs. Li + /Li. This makes the combination with TiS 2 unappealing due to the resulting low cell voltage of 1.6 V. The need for a high voltage lithium intercalation cathode material was met by the layered compound LiCoO 2 (LCO) which was developed by Goodenough

3 A BRIEF REVIEW: PAST, PRESENT AND FUTURE OF LITHIUM ION BATTERIES 1097 (а) (b) Co Li Co LiCoO2 Potential, V x in Li x CoO Fig. 2. (a) Rhombohedral structure of lithium cobalt oxide showing the layered nature, reprinted with permission from ref. [22], copyright 1980 Elsevier. (b) Voltage profile of LCO with LiPF 6 in ethylene carbonate / dimethyl carbonate (2 : 1), reprinted with permission from ref. [23], copyright 1996 the Electrochemical Society. et al. in 1980 [22]. The LCO structure is represented by a R( 3) mspace group with alternating layers of lithium oxide and cobalt oxide octahedrons (Fig. 2a). The large size difference of cobalt and lithium leads to a perfect layered material with little to no cation mixing (r Li = 0.76 Å, r Co = Å). This aspect makes the synthesis of stoichiometric and phase pure LCO very easy under various conditions. Lithium can be safely and reversibly removed up to 0.5 mol leading to a composition of Li 0.5 CoO 2 and an upper cut off voltage of 4.2 V vs. lithium (Fig. 2b) [23 25]. LCO has a poor thermal stability which worsen upon charging beyond 0.5 mol of lithium. Furthermore, structural distortions and cobalt dissolution are known to appear in the voltage region above 4.2 V vs. lithium [26, 27]. With the high voltage material LCO the only missing part was a stable, low voltage anode material which ideally could intercalate lithium in a similar fashion as TiS 2 does. In 1976 Besenhard et al. studied the electrochemical intercalation of Li + + and NR 4 into graphite and found intercalation potentials close to the thermodynamic potential of metallic lithium [28]. Using lithiated graphite instead of lithium as anode material seemed very intriguing at the time and a synthesis employing molten lithium was developed by Basu et al. in 1979 [29]. Combining LiC 6 and TiS 2 in a molten electrolyte gave rise to a reversible lithium intercalation battery which was patented in 1981 by the Bell labs [30]. The operating temperature was dependent on the employed salt or eutectic and ranged from 375 to 500 C. Clearly the use of the battery was somewhat limited due to the high temperature requirements but thousands of cycles were claimed. First attempts at cycling graphite versus lithium in non-aqueous liquid electrolytes were unsuccessful due to solvent co-intercalation into the lamellar structure of graphite [31 33]. The intercalation of solvated lithium leads to a substan- tial swelling and exfoliation of the material and a quick capacity fading is observed. By using a solid polymer electrolyte it was possible to demonstrate reversible lithium intercalation into graphite. The reported swelling of only 10% is the value theoretically expected for nonsolvated lithium intercalation [34]. A decisive discovery into the stability of graphite anodes in nonaqueous electrolyte solutions was made in 1990 by Dahn et al. [35]. It was shown that by using a solvent mixture of propylene carbonate (PC) and ethylene carbonate (EC) lithium can be reversible intercalated with little fading over the tested 19 cycles for graphite. This finding was attributed to the formation of a stable SEI during the first intercalation cycle at around 0.8 V vs. Li + /Li facilitated by the presence of EC. The SEI formation, nature and stability on graphite anodes was subject of extensive research in the 1990 s and is still of great interest today [36 42]. The SEI should be ideally formed before the actual lithium intercalation takes place. This makes graphite (Li + intercalation at <0.25 V [43]) favorable over amorphous carbons (Li + intercalation <1.5 V [44]). A comprehensive review on studies related to the SEI was given by Novák et al. [36]. Solvent mixtures of cyclic ester (e.g. ethylene carbonate (EC), propylene carbonate (PC)) and linear esters (e.g. dimethyl carbonate (DMC), diethyl carbonate (DEC)) with LiPF 6 are commonly used in commercial lithium ion batteries of today. A battery based on a LCO cathode, a petroleum coke anode and a non-aqueous electrolyte solution (PC:DEC; 1 M LiPF 6 ) was commercialized by Sony in 1991 and has been widely successful [45]. Since no metallic lithium is present in this configuration the name lithium ion battery (LIB) was chosen as a marketing term. Assembling the cell in a discharged state was a novel idea and has caught on in the lithium ion

4 1098 FLORIAN SCHIPPER, DORON AURBACH Current, ma Potential, V Fig. 3. Cyclic voltammorgam of LiMnO 2 showing the first 5 cycles at 10 μv s 1 from 2.0 to 4.6 V vs. Li + /Li; reprinted with permission from ref. [66], copyright 1999 the Royal Society of Chemistry. 5.0 battery community and is the standard for all commercial LIB nowadays. The stable SEI formation and absence of graphite exfoliation in EC containing electrolyte solutions lead to the manufacturing of commercial cells with a graphite anode from 1993 onwards. LITHIUM ION BATTERIES FROM COMMERCIALIZATION TO TODAY After the introduction of LIB to the consumer market the versatility of this new type of battery was soon adapted to portable devices. The high volumetric energy density of early LIB was around 200 W h L 1. Twice as high as competing systems at the time, namely nickel cadmium and nickel metal hydride batteries [46, 47]. This made LIB an ideal choice for laptop computers, digital cameras and cell phones which rapidly dropped in size and weight in 1990 s which can be partly attributed to LIB. LIB soon outperformed the competition and are now the leading battery technology for portable devices with a multi-billion dollar market and an overall market share of more than 60% worldwide [48, 49]. The workhorse amongst the different cathode types still is LCO which, owing to its high tap density of 2.5 g cm 3, can achieve high volumetric energy densities. The last two decades have seen an improvement in energy density from ~200 W h L 1 to values >400 W h L 1, which was mostly achieved by better cell engineering [47]. After the establishment of LIB in the battery market for portable consumer electronics the next big leap will be the development of adequate LIB for electric vehicles (EV). Even though LCO outmatches other lithium intercalation cathodes in terms of volumetric energy density the low capacity of 135 ma h g 1 leads to a low gravimetric energy density. High energy densities are necessary to achieve long driving ranges which are crucial to make full EV s appealing to the consumer market. The time before the first LIB and surely the last 25 years have seen the development of a number of different lithium intercalation cathodes [50]. The current chapter will deal with a selected number of those cathode materials which have seen commercialization or are under serious consideration. The very interesting work on high capacity lithium anodes (e.g. silicon [51, 52], tin oxide [53] etc.) and the latest development into new electrolyte solutions will be covered briefly at the end of the chapter. The interested reader is referred to the following articles for more detailed reviews on the subject of high capacity anodes and electrolyte solutions [54 59]. Layered Lithium Transition Metal Oxides Lithium Manganese Oxide LiMnO 2 In an attempt to switch to a cheaper and environmentally benign material LiMnO 2 was perused as an alternative to LCO. The work on LiMnO 2 began in the mid 1990 s with various failed attempts to synthesis the stoichiometric compound [60, 61]. Unfortunately, layered LiMnO 2 is unstable at elevated temperatures and an electrochemically inactive orthorhombic phase is thermodynamically favored [62]. It was independently established by Bruce and Delmas that an ion exchange of NaMnO 2 with lithium salts in organic solvents and temperatures below 150 C gives the layered LiMnO 2 compound [63, 64]. Another possible way to synthesis LiMnO 2 was realized via a low temperature hydrothermal (160 C, 3.5 days) approach [65]. LiMnO 2 is described by a monoclinic lattice with the C2/m space group. Electrochemical investigation of LiMnO 2 showed a large initial charge capacity of about 220 ma h g 1 but a significant drop to 130 ma h g 1 was observed for subsequent cycles [66]. The first cycle differs markedly from the following as was analyzed by cyclic voltammetry (Fig. 3). A peak around 3.6 V is observed which is lost for the following cycles. Furthermore, the peak at 4 V diminishes from the first to the fifth cycle. Once lithium is deintercalated to a stoichiometry of Li 0.5 MnO 2 a simple rearrangement of the lattice will lead to LiMn 2 O 4 [66, 67]. This is achieved by manganese ions moving to empty octahedral sites in the lithium layer. The rearrangement is quite significant and a large portion of the active material is already transformed to the LiMn 2 O 4 spinel after only 5 cycles [67]. The high initial capacity loss and the spinel formation render the otherwise interesting material useless for commercial application. In order to overcome the structural instability of LiMnO 2 a partial substitution of Mn 3+ with cobalt was pursued by Bruce et al. [68, 69]. A replacement of 10% of manganese for cobalt was reported to be enough to prevent the initial

5 A BRIEF REVIEW: PAST, PRESENT AND FUTURE OF LITHIUM ION BATTERIES (а) 5 (b) E, V 2 E, V Q, ma h g 1 (c) Q, ma h g 1 (d) E, V 2 E, V Q, ma h g Q, ma h g Fig. 4. Voltage profiles of LiNiO 2 derived under different synthesis conditions; reprinted with permission from ref. [70], copyright 1993 Electrochemical Society: (a) LiNO 3 + NiCO 3 in O 2 (750 C), R( 3 )m, I(003)/I(104) = 1.32; (b) LiNO 3 + NiCO 3 in Air (750 C), R( 3 )m, I(003)/I(104) = 1.25; (c) LiNO 3 + Ni(OH) 2 in O 2 (750 C), R( 3 )m, I(003)/I(104) = 1.39; (d) Li 2 CO 3 + NiCO 3 in O 2 (750 C), mix of R( 3 )m and Fm( 3 )m, I(003)/I(104) < 0.6. capacity loss. The LiMn 0.9 Co 0.1 O 2 material which now possess a R( 3) m space group still demonstrates a spinel formation which can be suppressed if higher cobalt contents are chosen. Lithium Nickel Oxide LiNiO 2 The nickel analogous to LCO was studied as a cheaper alternative to LCO. LiNiO 2 is isostructural to LCO with an R( 3) m space group and a theoretical capacity of 273 ma h g 1. However, LiNiO 2 is difficult to synthesize in a pure state due to the instability of the trivalent nickel at high temperatures [70]. The nonstoichiometry results in a displacement of nickel and lithium. Even small amounts of nickel within the lithium layer are known to reduce the discharge capacity and cycleability considerably [70, 71]. If the nickel content in the lithium layer reaches 50% a new cubic LiNiO 2 phase with the Fm( 3) m space group is formed which is electrochemically inactive. A qualitative measure to determine the magnitude of the nickel and lithium mixing is given by the intensity ratio of the X-ray reflections I(003)/I(104). The (003) reflection is prominent in R( 3) m but is absent in the Fm( 3) m modification of LiNiO 2 [72]. The synthesis temperature, synthesis gas, precursors as well as the stoichiometry have to be controlled very carefully in order to gain electrochemically active lithium nickel oxide [70, 73] (Fig. 4). The more complicated synthesis and poor thermal stability makes LiNiO 2 less attractive for LIB. Even though full cells with a carbon anode were investigated they never saw commercialization [74, 75].

6 1100 Safety LiMnO 2 (а) FLORIAN SCHIPPER, DORON AURBACH Rate LiCoO LiCoO 2 LiNiO 2 LiMnO LiNi 1/3 Co 1/3 Mn 1/3 O 2 E E E Co 3+/4+ :e g Mn 3+/4+ :e g Co 3+/4+ :t Ni 3+/4+ :e g 2g 0.8 Ni 3+/4+ :t Mn 3+/4+ :t 0.2 Capacity Co 2 2g 2g :2p Co 2 :2p Co 2 :2p N(E) N(E) N(E) LiNiO 2 LiNi 0.5 Mn 0.5 O 2 >180 ma h g ma h g ma <150 ma h g 1 h g 1 (b) Fig. 5. (a) Phase diagrams of layered lithium NCM with the end members LiNiO 2, LiMnO 2 and LiCoO 2 ; reprinted with permission from ref. [87], copyright 2015 WILEY-VCH. (b) Electronic structure of the individual endmembers of NCM: LiCoO 2, LiNiO 2 and LiMnO 2, reprinted with permission form ref. [88], copyright 2008 the Royal Society of Chemistry. Partially substitution of nickel with manganese was studied for Li x Ni (1 y) Mn y O 2 with y up to 0.6 [76]. Since pure lithium manganese oxide is not isostructural to LiNiO 2 higher manganese ratios could not be realized. Furthermore, the measured capacity decreased with increasing manganese content which made the material uninteresting for further research at first. The material was revisited a decade later with greater success, it appears that the annealing temperature has to be above 800 C to yield an electrochemically active material [77, 78]. The composition LiNi 0.5 Mn 0.5 O 2 is of great interest and has sparked a lot of research. It was discussed to consist either of Ni 3+ /Mn 3+ or Ni 2+ /Mn 4+. The high capacity (>200 ma h g 1 ) suggested the latter which was confirmed by in-situ X-ray absorption studies and by first principle calculations [78, 79]. Problems are arising from the synthesis and ca. 10% of Ni 2+ can be found in the lithium layer impeding the lithium intercalation [80, 81]. Materials prepared by ion exchange from NaNi 0.5 Mn 0.5 O 2 show less cation mixing and good rate capabilities were reported suggesting that LiNi 0.5 Mn 0.5 O 2 is indeed a worthwhile endeavor [82]. The substitution of nickel with cobalt improves the cycling behavior tremendously when the cobalt content is 0.2 in Li x Ni (1 y) Co y O 2 [74]. This fact was attributed to a better structural stability of the cobalt doped material and stoichiometric compounds are easily prepared in stark contrast to pure LiNiO 2 [83, 84]. With more cobalt a decrease of nickel/lithium mixing was observed which explains the enhanced electrochemical performance of the material [85]. Lithium Nickel Cobalt Manganese Oxide Li[Ni 1 x y Co x Mn y ]O 2 The work on doping of the individual endmembers of LiCoO 2, LiNiO 2 and LiMnO 2 eventually culminated in the development of intercalation compounds based on all three transition metals with the general formula Li[Ni 1 x y Co x Mn y ]O 2 [86]. The notation most common in literature is the abbreviation NCM or NMC with numbers indicating the decimals of the transition metals (e.g. NCM523 = Li[Ni 0.5 Co 0.2 Mn 0.3 ]O 2 etc.). The different traits associated with manganese, cobalt and nickel can be summarized as safety, rate/capacity retention and capacity, respectively (Fig. 5a). Manganese is electrochemically inactive in the NCM material and can therefore stabilize the structure and acts as environmentally friendly and inexpensive filler (Fig. 5b). Cobalt is necessary to facilitate an easy synthesis with as little structural and stoichiometric impurities as possible (e.g. no cation mixing). But cobalt is also prone to react with the electrolyte if the Co 3+/4+ redox region is reached which is the case for >0.5 mol of lithium extraction from LiCoO 2. The Ni 3+/4+ redox e g band shows only a small overlap with the O 2 2p band and much more lithium can be extracted from LiNiO 2 in theory. In order to achieve a viable cathode material the transition metal ratio has to be controlled under aspects of high capacity (more nickel) or better cycle stability (more cobalt) and safety (thermal stability)/cost (manganese). The symmetrical composition NCM333 was studied by Ohzuku et al. and was found to be cycle stable with a capacity of ca. 150 ma h g 1 when cycled between V [89]. Extending the voltage window up to 4.6 V will increase the capacity to values above 200 ma h g 1 but at the cost of cycle stability [90]. NCM333 is simple to synthesis via solid state, sol gel or co-precipitation reactions and can be calcined in air at 1000 C without any structural impact [90 92]. Since NCM333 delivers a higher capacity than LCO it is used for LIB in some of the current full and hybrid electric vehicles.

7 A BRIEF REVIEW: PAST, PRESENT AND FUTURE OF LITHIUM ION BATTERIES 1101 The urge to increase the capacity of the NCM family is usually met by increasing the nickel content to compositions >50% nickel which are referred to as nickel rich materials. The compositions NCM622 and NCM811 are under intense scrutiny and capacities of 160 and 190 ma h g 1 have been reported respectively [93, 94]. One problem with nickel rich NCM intercalation materials arises during the synthesis. An ambient air atmosphere is detrimental to the material capacity due to the easy formation of Li 2 CO 3 and LiOH on the materials surface. Hence, the synthesis has to be conducted under an oxygen atmosphere with low CO 2 and water levels. Even storage of the material is problematic due to the same reasons. A considerable increase in electrochemically inactive LiOH and Li 2 CO 3 has been reported to occur after only 7 days of ambient storage [95]. The extent to which this will occur is related to the cobalt content with less parasitic reactions with higher cobalt levels. Cobalt is believed to act as a structural stabilizer in the material. The formation of lithium hydroxide and lithium carbonate is accompanied by an oxygen loss in the material which is compensated by a reduction of Ni 3+ to Ni 2+. One way to compensate the surface reactivity is given by surface coating, e.g. Al 2 O 3 [95]. Other pursuits to overcome these issues can be found in gradient materials. Since the side reaction occur on the outer surface of the material a gradual reduction from a nickel-rich core to a nickel-poor shell leads to a more stable material (Fig. 6) [96]. The second drawback found in nickel rich materials is the low thermal stability. The onset temperature of NCM materials is depending on the cobalt content and reaches values of ca. 250 C for NCM433 and NCM523. With Ni > 50% a rapid decrease to 170 C (NCM622) and ca. 145 C (NCM811) is observed (all measured in charged state; 4.3 V vs. Li + /Li) [97]. Strategies to mitigate the temperature instability are ranging from surface coating with AlPO 4 [98], AlF 3 [99] to gradient materials [100]. It is clear that a high capacity material is attractive for EV s but considerable focus must be paid to its safety aspects. Surface = Mn = Ni Interior Fig. 6. SEM images and schematics of gradient nickel rich NCM materials with a nickel rich core and manganese rich shell, reprinted with permission from Ref. [101], copyright 2011 the American Chemical Society. Lithium Nickel Cobalt Aluminum Oxide Li[Ni 1 x y Co x Al y ]O 2 Pursuits to stabilize the LiNiO 2 structure went beyond the previously discussed manganese and cobalt doping. Further attempts included inactive Al 3+ as replacement for Ni 3+ with the structure Li(Ni 1 x Al x )O 2 [102, 103]. The material shows an initial capacity of 150 ma h g 1 which is smaller than the usual 200 ma h g 1 found for LiNiO 2. A short cycling test with 30 cycles renders pure LiNiO 2 at ca. 100 ma h g 1 while Li(Ni 3/4 Al 1/4 )O 2 still delivers ca. 140 ma h g 1 [102]. The lower capacity is not surprising since Al 3+ does not partake in the electrochemistry of the material and hence limits the lithium removal. Aluminum acts as a structural stabilizer within the nickel layer and no phase changes occur during cycling leading to an enhanced cycle stability and greatly increased thermal stability [102, 104]. The ion radii for Ni 3+ and Al 3+ on octahedral sites are 0.56 and Å respectively and it is not unexpected that Li(Ni 1 x Al x )O 2 shows a cation mixing with Al 3+ and Ni 3+ in the lithium layer [104]. To get the best of both worlds a well ordered layered LiNiO 2 with little cation mixing via cobalt doping and a good thermal stability via aluminum doping, the Li[Ni 1 x y Co x Al y ]O 2 (NCA) material was developed [105]. The popular composition Li[Ni 0.8 Co 0.15 Al 0.05 ]O 2 was shown to have a better cycle stability and less impedance growth in a full cell setup compared to Li[Ni 0.8 Co 0.2 ]O 2 [106]. Quite recently the Li[Ni 0.81 Co 0.10 Al 0.09 ]O 2 composition was proposed as a viable NCA candidate. A good electrochemically performance and improved thermal stability over the classic Li[Ni 0.8 Co 0.15 Al 0.05 ]O 2 was demonstrated [107]. Lithium Manganese Spinel LiMn 2 O 4 Manganese based intercalation materials are quite attractive due to the low cost and low environmental impact of manganese. LiMn 2 O 4 (LMO) was first proposed by Thackeray in 1983 and subsequently extensively studied by his group [108, 109]. The spinel configuration is very attractive in terms of lithium intercalation. LMO shows a Fd3m space group with oxygen forming a cubic closed package while lithium occupies 1/8 of the tetrahedral (8a) sites and manganese 1/2 of the octahedral (16d) sites (Fig. 7a) [109]. The empty octahedral (16a) sites share edges with the 8a site and form a three dimensional pathway for lithium ion intercalation and deintercalation (Fig. 7b). LMO

8 1102 FLORIAN SCHIPPER, DORON AURBACH LiO 4 MnO 6 Е, V ~ 0 (а) (c) Li x Mn 2 O (I) (II) (III) 50 (b) Q, ma h g 1 8a site (Li + ) c site (vacancy) Fig. 7. (a) Crystal structure of LiMn 2 O 4 spinel with lithium sites in orange and manganese sites in blue adapted with permission from ref. [113], copyright 2012 the Royal Society of Chemistry; (b) 3D pathway for lithium ion diffusion by linked octahedral vacancies and lithium sites adapted with permission from ref. [113], copyright 2012 the Royal Society of Chemistry; (c) typical discharge profile for Li x Mn 2 O 4 from x = 0.3 to x = 1.8 ( V) reprinted with permission from ref. [111], copyright 1990 The Electrochemical Society. demonstrates two voltage regimes with two voltage plateaus around 4 V and a voltage plateau at ca. 3 V (Fig. 7c) [110]. The insertion/extraction of lithium ions in the 8a sites takes place at the 4 V region while preserving the initial cubic spinel lattice but two phases can be identified [111]. The 3 V plateau relates to an insertion of lithium ions into the vacant 16c sites accompanied by a change to a tetragonal phase driven by the Jahn Teller distortion of the Mn 3+ in Li 2 Mn 2 O 4. [111, 112]. The phase change is accompanied by a 5.6% increase in unit cell volume which is detrimental to the structural integrity of the material [111]. In addition, the voltage drop of 1 V is considerable and might not be suitable for electronics. Operating LMO within the 4 V region gives a capacity of ca. 120 ma h g 1 and leads to a composition in the delithiated state of Li 0.2 Mn 2 O 4. Even if the material is operated solely in the upper 4 V region a rapid capacity fading is observed which becomes intolerable at elevated temperatures [ ]. Several mechanisms are used to describe the poor cycle stability of LMO with manganese dissolution being the most prominent. The disproportionation of Mn 3+ leads to the formation of Mn 4+ and Mn 2+ the latter of which is readily dissolved into the electrolyte. The manganese dissolution is an inherent problem with manganese containing active materials and is more pronounced at elevated temperatures and voltages [ ]. Especially conditions which favor the electrolyte decomposition to HF are known to accelerate the manganese dissolution [118]. Since the dissolution of manganese originates from the Mn 3+ ion efforts to increase the average oxidation state of manganese above 3.5 were undertaken. Replacing a portion of manganese with Li +, Mg 2+, Zn 2+ will increase the overall oxidation state of manganese and was reported to enhance the cycle stability [119]. However, this increase in stability comes at the cost of capacity and stable capacities of ca. 100 ma h g 1 were achieved. Since the Mn 3+/4+ redox couple is the active species in LMO an increase in oxidation state of manganese must be accompanied by a reduced capacity if the lower voltage plateau is avoided. Replacing manganese with Al 3+ on the other hand will not affect the oxidation state and was shown to enhance the capacity retention even at 45 C [120, 121]. The replacement of manganese by aluminum leads to a lower a parameter in the cubic lattice due to the smaller size of the Al 3+ ions compared to Mn 3+ [120]. The composition of LiMn 1.7 Al 0.3 O 4 was found to be optimal in terms of capacity retention with a capacity of ca. 110 ma h g 1. Increasing the Al content further will gradually decrease the capacity since more and more Mn 3+ is replaced by electrochemically inactive Al 3+. The better stability was attributed to a better structural integrity for the aluminum doped LMO which leads to a more stable unit cell parameter a over 50 cycles [120]. Another approach to enhance the cycle stability can be found in various coatings of the active material. The clear aim here is to minimize the contact area between the material and the electrolyte with electrochemically inert materials (e.g. MgO [122], CeO 2 [123] ZnO [124], AlF 3 [125], AlPO 4 [126], LiAlO 2 [127] etc.). Most commonly the coatings are believed to be a HF scavenger which in turn would lead to suppression in manganese dissolution. One general problem with wet chemical coatings is the non-uniform coating of usually nano- and sub micrometer sized particles onto the active material which leaves a lot of surface area exposed to the electrolyte. Quite recently powder atomic layer deposition has become popular in various research fields and its viability in case of LiCoO 2 was already demonstrated [128]. Revisiting some coatings for LMO with a powder ALD process could be an interesting study since a uniform thin layer might be more adequate in shielding the active material form the electrolyte. Despite its shortcomings LMO is commercially used in power tools were a high power output is

9 A BRIEF REVIEW: PAST, PRESENT AND FUTURE OF LITHIUM ION BATTERIES Ni 4+ Ni 3+ Ni 2+ Potential, V vs. Li + /Li Discharge capacity, ma h g Rate, h x in Li x Ni 0.4 Mn 1.6 O 4 Mn 4+ Mn Fig. 8. Voltage profile for Li x Ni 0.4 Mn 1.6 O 4 with rate capability test (inset), reprinted with permission from ref. [136], copyright 2009 Elsevier. required which can be met due to the 3D intercalation pathway. The significant reduction in cost makes blended electrodes of LMO and NCM attractive for EV s Furthermore, better storage capability [129], enhanced thermal stability [130] and better capacity retention [131] haven been demonstrated for LMO/NCM blends. Lithium Nickel Manganese Spinel LiNi 0.5 Mn 1.5 O 4 An interesting offspring of the work on metal doped LMO material (e.g. Ni, Cr, Co) are high voltage spinels which were missed in the first instance. Several works into LiM x Mn 0 x O 4 (M = Ni, Cr, Co) type materials found an increased cycle stability but at the cost of capacity depending on the amount of dopant [110, 132]. Later studies showed an additional voltage region around 5 V which was initially missed and gave rise to the name 5 V spinels [133, 134]. The length of the 5 V region is dependent on the amount of dopant and LiNi 0.5 Mn 1.5 O 4 (LNMO) arose as a promising material with a theoretical capacity of 147 ma h g 1 (Fig. 8) [135]. The average oxidation state of manganese is Mn 4+ in LNMO and is unchanged during cycling with nickel being oxidized from Ni 2+ to Ni 3+ and finally to Ni 4+. Since manganese is pinned in the +IV oxidation state Jahn Teller distortions of the Mn 3+ cation are avoided and the manganese dissolution should be negligible. Indeed, samples stored at 55 C for 7 days showed a Mn dissolution of 0.3% for LNMO while LMO showed 3.2% [137]. Nonetheless, even the lower dissolution still leads to a capacity fading during cycle and higher manganese dissolutions are observed in cycled cells [138, 139]. The high voltage plateau of ca. 4.7 V vs. Li + /Li for LNMO is beyond the stability of current electrolyte solutions Tests monitoring the manganese dissolution based on the state of charge (100% SOC refers to the upper cut off voltage in the delithiated state) showed an increased Mn dissolution with an increase in SOC and temperature [140]. Furthermore, the formation of NiF 2 and MnF 2 on the cathode surface was shown due to the presence of trace HF in the electrolyte solution. The synthesis of LNMO is a problematic part and the material can be found in two space groups namely the ordered spinel P or the disordered Fd ( 3) m [136]. The latter was shown in multiple studies to outperform the ordered phase in terms of rate capability and cycle life [ ]. Amatucci et al. argued that the higher performance in the Fd ( 3) m phase is due to its higher lithium diffusivity and similar findings were made by Sun et al. for titanium substituted LNMO [142, 144]. Ways to shift the synthesis towards the disordered phase are cation substitution of nickel and manganese with different metals (e.g. Fe, Co, Ti) [142, 143], the synthesis of oxygen non-stoichiometric compounds LiNi 0.5 Mn 1.5 O 4 δ by adjusting the calcination temperature [141, 145] or by synthesizing nickel deficient (LiNi 0.4 Mn 1.6 O 4 ) material [136]. Furthermore, Li x Ni 1 x O and NiO impurities are likely to

10 1104 FLORIAN SCHIPPER, DORON AURBACH (a) (b) Li FeO 6 PO 4 Capacity, ma h/g ma/cm 2, 2.1 ma/g 4 a c Cell voltage, V 3 Charging 1st Cycle 2 5th Cycle 15th Cycle 25th Cycle у in Li 1 y FePo 4 Fig. 9. (a) LiFePO 4 structure with 1D lithium channels, reprinted with permission from ref. [156], copyright 2000 Elsevier. (b) Typical voltage profile for LiFePO 4 from 3.9 to 2 V, reprinted with permission from ref. [152], copyright 1997 The Electrochemical Society. form during the synthesis. Close attention has to be paid to synthesis conditions such as O 2 pressure, calcination temperature and nickel content in order to keep the impurities to a minimum [133, 143, 146]. The stability of LNMO in electrolyte solution during the high voltage cycling is a problematic aspect. Ways to shield the material were investigated mainly in form of surface coatings [ ]. Liu and Manthiram recently showed some promising results with selfassembled surface layers of Al 2 O 3, ZnO, Bi 2 O 3 and AlPO 4. They form a shell like coating around the LNMO particles and enhanced capacity retentions were reported [150]. Lithium Iron Phosphate LiFePO 4 The olivine LiFePO 4 (LFP) was identified as an intercalation compound by Goodenoughs group in 1997 (Figs. 9a, 9b) [151, 152]. Initial research demonstrated only small capacities of less than 120 ma h g 1 but the material gained considerable research interest due to the cheap and benign nature of iron compared to cobalt and nickel. Furthermore, LFP demonstrates a flat voltage profile with a single plateau at ca V vs. Li + /Li which makes the material more stable in non-aqueous electrolyte solutions. Unlike LCO, the Fe 2+/3+ redox couple of LFP does not react with stateof-the-art electrolyte solutions even in a charged state. In addition the thermal stability of LFP is excellent with no exothermic reactions up to 400 C (powder measurement) [153]. The problems associated with LFP are the low electrical conductivity of < 10 9 S cm 1 [154] and the poor lithium diffusivity with a diffusion coefficient D Li of to cm 2 s 1 (from LiFePO 4 to FePO 4 ) [155]. A heads on approach to overcome the low electrical conductivity is given by various carbon coating techniques. An intimate contact between LFP and a conductive carbon layer leads to an increased electrical conductivity. Amongst others resorcinol formaldehyde gels [157], ball milling [158] and sol gel process [159] have been reported to be viable options to achieve increased performance of LFP. In 2002 Chiang et al. reported on LFP with an 8 orders of magnitude higher electrical conductivity [160]. The authors used different dopants from Mg, Ti, Zr to Nb with the highest electrical conductivity for a 1 at % Nb doping in exchange for lithium in LFP. The tremendous conductivity increase was attributed to a p-type semiconductor behavior in the doped LFP. A later study by Nazar et al. suggested the effect originates from a percolating nano-metallic iron phosphide network [161]. The olivine structure of LFP consist of LiO 6 octahedra, FeO 6 octahedra and PO 4 tetrahedra with only one dimensional lithium ion channels for intercalation (Fig. 9a). The limited diffusion pathways explain the low D Li which in combination with micron size particles leads to a poor electrochemical performance of the material. Nano-sized LFP was shown to approach the theoretical capacity of 170 ma h g 1 at room temperature and a C/10 rate with a decent rate capability [157, 162, 163]. LFP benefited greatly from material engineering and nanotechnology which allows to minimize the impact of the low diffusion coefficient by shortening diffusion lengths. One drawback of nanosized materials is the lower tap density compared to their micron sized equivalent, hence nano-lfp

11 A BRIEF REVIEW: PAST, PRESENT AND FUTURE OF LITHIUM ION BATTERIES 1105 Potential (Li 2 MnO 3 vs. Li), V Charge (a) Discharge С 55 С Capacity, ma h g 1 (charge removed/inserted per gram of Li 2 MnO 3 ) Potential, V vs. Li/Li C 900 C 1000 C 900 C 850 C (b) 800 C 850 C Capacity, ma h/g 800 C 600 C 400 C 600 C 400 C 10 ma/g first cycle Fig. 10. (a) First charge/discharge cycle of Li 2 MnO 3 at 30 and 55 C; reprinted with permission from ref. [183], copyright 2003 the American Chemical Society; (b) First charge/discharge cycle of nano-li 2 MnO 3 at 25 C at various synthesis temperatures with the 600 and 400 C samples having stacking faults, particle size is increasing with synthesis temperature; reprinted with permission from ref. [191], copyright 2009 The Electrochemical Society. demonstrates a lower volumetric energy density. It is noteworthy that other intercalation cathodes can benefit from nanotechnology as well to a certain extent. Total capacity and rate capability are usually higher in nano-sized materials but so is their reactive surface area. This is especially problematic if the material tends to react with the electrolyte, like LCO. Combined approaches with nano-sized and electrical conductive LFP makes the material a viable choice for LIB and thousands of cycles were demonstrated in lab scale experiments [164, 165]. The lower operating voltage of 3 to 4 V allows for a long cycle life (no side reactions with the electrolyte) and the only limiting factor in full cells will be the graphite anode. Since long cycle life, thermal stability (safety), cost and environmental impact are of great interest for large battery applications, LFP is used in some current generation EV s. Furthermore, combinations with the high voltage lithium titanate spinel anode (Li 4 Ti 5 O 12, 1.5 V vs. Li + /Li, 175 ma h g 1 [166]) show great promise for stationary applications were very long cycle life is a key aspect and even up to cycles have been demonstrated at a 10C rate [ ]. The olivine family has been expanded over the years by LiMnPO 4 [170, 171], LiCoPO 4 [172] and LiNiPO 4 [173] which show intercalation voltages of 4.1, 4.8 and 5.1 V respectively. The higher voltages seem appealing at a first glance to increase the energy density but are out of the stability window for current electrolytes. So far no useful capacities have been demonstrated for LiCoPO 4 and LiNiPO 4. LiMnPO 4 (LMP) seems to be a good candidate to replace LFP and increase the energy density along the way. Unfortunately, LiMnPO 4 shows an electrical conductivity 3 orders of magnitude lower than LFP and large polarizations are usually observed [174]. It was shown that the polarization effect can be greatly reduced by nano-sizing and carbon coating LMP yielding a capacity of ca. 100 ma h g 1 with a voltage plateau at ca. 4.1 V [174] and approaching 150 ma h g 1 at a C/20 rate [175]. Another way to increase the intercalation voltage of LFP can be found in replacing some of the iron with manganese leading to a solid solution of LFP and LMP. Even though tests demonstrate a better capacity if compared to LMP and a higher average voltage compared to LFP the voltage profile is split in two regimes [176, 177]. Lithium intercalation occurs at 4.1 and 3.5 V corresponding to the individual materials. It is doubtful that such material will find its way to commercialization due to the high voltage drop of ca. 600 mv. Improvements in high voltage electrolytes could be very interesting and would allow to study the high voltage olivines LiCoPO 4 and LiNiPO 4 adequately. Reported total capacities and capacity fading for these materials are unsatisfactory and fall short when compared to LFP. It is unclear if the electrochemical behavior is intrinsic for the material or just reflects the breakdown of the electrolyte solutions [178, 179]. The materials might benefit from recent research into room temperature ionic liquids used as 5 V electrolyte solutions with 0.5 M LiTFSI with a LiNi 0.5 Mn 1.5 O 4 cathode [180]. Lithium Rich Layered Materials xli 2 MnO 3 (1 x)li(ni 1 y z Co y Mn z )O 2 The layered Li 2 MnO 3 (space group C2/m [181, 182]) gained considerable research interest in the early 2000 s due to its anomalous high charge capacity of 320 ma h g 1 when charged to 4.7 V at 55 C (Fig. 10) [183]. The +IV oxidation state of manganese in Li 2 MnO 3 would not allow for any lithium extraction even though some authors claimed the formation of

12 1106 FLORIAN SCHIPPER, DORON AURBACH Voltage, V vs. Li/Li Capacity, ma h g ma h g ma h g Fig. 11. Voltage fade of lithium rich NCM over 36 cycles; reprinted with permission from ref. [200], copyright 2013 Elsevier. Mn 5+. Careful XPS and redox titration studies showed no indication of Mn 5+ in the charged material [184, 185]. Other authors suggested a mechanism involving the release of oxygen from the lattice forming Li 2 O and MnO 2 on the material surface for Mn 4+ species [186, 187]. A third explanation was given by Bruce et al. based on the oxidation of the electrolyte under the high voltage condition according to the simplified reaction: R H R 0 + H + + e. Hence, providing electrons (charge) and protons, the latter of which are exchange with lithium ions of the active material (H + Li + ) [183]. The breakdown of the electrolyte is a complicated matter in itself and gas evolution including CO 2 and oxygen are known to occur on overcharge [188, 189]. The Li 2 MnO 3 charge mechanism is still not fully understood and is part of an ongoing debate in the field. Authors usually favor a combined explanation involving oxygen release and proton exchange [190, 191]. Anyhow, L 2 MnO 3 can deliver a high charge capacity on the first charge at elevated temperatures but fails to retain this capacity already on the first discharge. Attempts to obtain a high capacity at room temperature and enhanced cycle stability were pursuit by nano-sizing the material (Fig. 10b) [192, 193]. It was claimed that defects are leading to an increase in capacity. The defect structures are derived at lower temperatures and the particle size is considerably lower in the defect structures. It is unclear if the increase is related to the particle size or indeed defects [191]. Even though nano-li 2 MnO 3 shows a much better performance with a first discharge capacity of 236 ma h g 1 a rather fast fading is observed with only ca. 190 ma h g 1 after 30 cycles [192]. The work on Li 2 MnO 3 sparked a rush of work on combination of layered materials with the lithium rich phase of the general formula xli 2 MnO 3 (1 x)limo 2. The ternary NCM system is under intensive scrutiny by a multitude of research groups and companies. Thackeray et al. studied the structure of the lithium rich NCM333 compounds with 5 and 10% of lithium excess (Li 1.05 Ni 1/3 Co 1/3 Mn 1/3 O 2 and Li 1.10 Ni 1/3 Co 1/3 Mn 1/3 O 2 ). It was confirmed by HRTEM convergent beam electron diffraction that two phases ( R( 3) m of NCM and C2/m of Li 2 MnO 3 ), integrated on a nanoscopic scale, are present [194]. The Li 2 MnO 3 phase is inactive if cycled to a low cutoff voltage of 4.4 V but becomes active once the voltage is increased to 4.6 V [195]. The capacities reported on lithium rich NCM materials are ranging from 200 up to 300 ma h g 1. Usually many factors are different starting with the synthesis conditions, voltage range used, cycling temperature and more importantly the values of x, y and z in xli 2 MnO 3 (1 x)li(ni 1 y z Co y Mn z )O 2 [ ]. A popular composition in this field seems to be the xli 2 MnO 3 (1 x)li(ni 1/3 Co 1/3 Mn 1/3 )O 2 with x = 0.3, 0.5 and 0.7. Drawbacks related to lithium rich materials are the faster capacity fading compared to their lithium neutral layered NCM counterpart and a continuous voltage fade with cycle progression (Fig. 11) [200]. The latter is often ascribed to a layered to spinel transformation and the former to a transition metal dissolution [201]. Naturally the first response was to coat the material with a variety of inactive compounds known from the LiMn 2 O 4 spinel research field (e.g. TiO 2 [202], FePO 4 [203], ZrO 2 [204], AlF 3 [205, 206] etc.). Another approach is found in rising the average oxidation state of manganese in the lithium rich composition. This can be achieved by either moving to a cobalt free composition (e.g. 0.3Li 2 MnO 3 0.7LiMn 0.5 Ni 0.5 O 2 [207]) or by applying a cobalt gradient with a manganese rich shell [208]. The research surrounding lithium rich materials is very active often aiming at increasing the capacity, capacity stability and suppressing the voltage fade. But usually without gaining decisive knowledge as to why and how all those parameters are linked and might influence each other. The integrated xli 2 MnO 3 (1 x)li(nicomn)o 2 material is a complicated subject and fundamental understanding is difficult to achieve with such a compound having multiple redox centers. Recently, Tarascon et al. set out to study model materials of the formula Li 2 Ru 1 x Sn x O 3. They, show no voltage decay but high capacities of ca. 280 ma h g 1 and have, with ruthenium, only one active redox center [209, 210]. A mechanism explaining the capacity was formulated to combine cationic oxidation as well as the formation of anionic peroxo-superoxo (2O 2 (O 2 ) n ) species [210]. A following study on Li 2 Ru 0.75 Ti 0.25 O 3 with a capacity of ca. 240 ma h g 1, but now with voltage fading, concluded the origin to be linked to the trapping of Ti 4+ cations in tetrahedral sites [211]. The

13 A BRIEF REVIEW: PAST, PRESENT AND FUTURE OF LITHIUM ION BATTERIES 1107 Table 1. Properties of a selected number of common solvents in LIB taken from ref. [58] Solvent Structure T m, C T b, C * Measured at 40 C. η, mpa s (25 C) ε (25 C) ρ, g cm 3 (25 C) O PC O O O EC O * O O DMC O O O EMC O O DME O O above mentioned work on fundamental understandings related to lithium rich materials was recently combined into a review paper and the interested reader is referred to ref. [212]. Lithium rich NCM materials are definitely an interesting and hot topic and are believed to be the successor of traditional NCM cathodes owing to their much higher capacity. Nonetheless, further research is necessary and fundamental science seems only now to catch up with the material development and some much needed insights into the material are gained. Understanding the material on a fundamental level will help to develop smarter compositions of Ni-rich materials and hopefully aid in their development towards commercialization. Electrolyte Solutions for Lithium Ion Batteries Electrolyte solutions for LIB are complicated systems consisting of a solvent mixture, a salt and any number of additives. Even though they are a vital part of any given battery technology, this topic may be less attractive for many researchers. Because work on electrolyte solutions and related interfacial problems may be less conclusive than work on cathode materials. The latter involves more bulk structural analysis that can provide more unambiguous results. The electrolyte solutions used in LIB of today consist of ethylene carbonate (EC) and linear carbonates (e.g. ethyl methyl carbonate (EMC), dimethyl carbonate (DMC)) as solvents in combination with LiPF 6 as a lithium salt. These solutions provide a reasonably wide electrochemical window for 4 5 V Li batteries, thanks to complicated passivation phenomena. Additives used in commercial electrolyte solutions are usually not disclosed. EC is a vital part and has to be present to facilitate the SEI formation to prevent graphite exfoliation as was described in the first part of this paper. Before LIB had become a commercial reality, research on electrolyte solutions was focused on metallic lithium and finding ways to prevent dendritic growth. The focus changed drastically with the use of high voltage intercalation cathode materials and with better understanding of lithium intercalation into soft carbons and later graphite. In general, a number of attributes are necessary for an electrolyte solution to be useable in LIB: (1) Wide operating temperature range (low meting point, high boiling point); (2) Low viscosity of the solvent mixture; promotes ionic conductivity; (3) Decent solubility and dissociation of the used lithium salt in the solvent mixture (high relative permittivity); promotes ionic conductivity; (4) Large electrochemical window to operate with low voltage anodes and high voltage cathodes. Set aside from the above points, other characteristics such as non-toxicity, environmentally benignity, thermal stability, and of course, low cost are sought after as well. Electrolyte solutions make up ca. 45% of a commercial LIB (by mass) and roughly 1/5-th of the total price generating a considerable market around them [213, 214]. Solvents The most used solvents in LIB are cyclic and acyclic carbonates but others have been studied as well (Table 1). The combination of solvents with lithiated graphite (LiC 6 ) or metallic lithium prevents the use of protic solvents. Only polar-aprotic solvents which dissolve an adequate amount of lithium salt and form

14 1108 FLORIAN SCHIPPER, DORON AURBACH Table 2. Properties of different lithium salts used in LIB, taken from ref. [58, 244] Salt Structure Al corrosion σ, ms cm 1 1 M in PC (25 C) σ, ms cm 1 1 M in EC DMC (25 C) LiPF 6 Li + [PF 6 ] No LiClO 4 Li + [ClO 4 ] No LiAsF 6 Li + [AsF 6 ] No O O O O LiBOB B Li + No ~5.0* ~11.0* O O O O LiTFSI Li + [N(SO 2 CF 3 ) 2 ] Yes * Measured with 0.7 M in PC or EC/DMC blends with esters [244]. solutions with high ionic conductivity are relevant. These requirements limit the spectrum of relevant solvents families to esters, ethers, alkyl carbonates, nitriles and sulfones. The combination of low viscosity solvents with viscous ones which have a high dielectric constant is an attractive choice and has become the common practical approach. Indeed most practical electrolyte solutions are based on EC and linear carbonates of lower viscosity. The high melting point of EC is a problem, but combinations with EMC or DMC lead to solvent mixtures with an acceptable liquid range and melting points below 20 C [58, 215]. The combinations of EC with solvents such as DMC, DEC, EMC and LiPF 6 became the standard solutions in LIB. Their improvement is an ongoing area of interest. Especially their thermal stability and questionable high voltage stability are considered as major problems that have to be addressed. Fluorinated solvents have been reported to improve on both matters [ ]. Fluorinated solvents may also act as flame retardants in LIB but should be used mostly as co-solvents due to their higher viscosity compared to their non-fluorinated counterparts [219]. Even though the C F bond is polar and adds to the overall relative permittivity of these solvent, their higher viscosity leads to lower ionic conductivity [220]. Fluorinated solvents have been demonstrated to improve the passivation properties of surface films formed on graphite and silicon anodes [217, 221, 222]. Salts A wide variety of salts was used throughout the development of LIB until a compromise was found with LiPF 6 (Table 2) [ ]. Any salt used in LIB has to be beneficial for passivating surface film formation, inert towards the current collector and thermally stable. LiClO 4 was shown to have a high solubility in aprotic solvents (high ionic conductivity) but forms less stable passivation layers on aluminum and most important, it may be explosive [226, 227]. LiAsF 6 was also tested intensively. However it is toxic, what rules out its commercial use but not necessarily its application in fundamental research. LiPF 6 has become the standard salt for LIB due to its good balance of properties such as high solubility and ionic conductivity, stable SEI formation in combination with EC and reasonable thermal stability [58]. One major drawback found in LiPF 6 is the HF formation with trace amounts of water according to [228, 229]: LiPF 6 LiF + PF 5, PF 5 + H 2 O POF 3 + 2HF. The as formed HF dissolves transition metal cations via exchange of protons by transition metal ions in the cathode structure. Thus, providing detrimental situations in which transition metal cations migrate to the anode side and are being deposited on the graphite surface. Their reaction on the graphite side worsens the anode passivation, leading to side reactions and consequently, lowering their capacity and increase their impedance [230, 231]. Transition metals can also be deposited as oxides or fluorides directly onto the surface of the cathode thus also leading to an increase in impedance and eventually cell death [232]. As with all chemical reactions, an increase in temperature enhances the rate of side reactions, thus accelerating the capacity degradation. LiPF 6 is far from being the perfect salt for LIB but makes for a decent compromise of properties. LiBOB (Li-Bioxalato-borate) has been under scrutiny since the late 1990 early 2000 and has attracted some commercial interest [233, 234]. The electrochemical window of electrolyte solutions employing LiBOB was reported to be >4.5 V making it a potential candidate for high voltage spinel cathodes or olivines [233]. The solubility is in general lower than

15 A BRIEF REVIEW: PAST, PRESENT AND FUTURE OF LITHIUM ION BATTERIES 1109 that for LiPF 6 leading to a lower ionic conductivity [235]. Some research demonstrated a better high temperature cycle stability for LiBOB since the salt has a good stability against trace amounts of water and does not form detrimental HF [236, 237]. A study by Dahn et al. focused on the thermal stability of LiBOB containing electrolyte solutions and showed a very good stability in combination with lithiated graphite. Unfortunately, the reactivity with Li 0.5 CoO 2 was higher than that for a standard solution of 1 M LiPF 6 in EC/DEC [238]. LiTFSI (Bis(trifluoromethane) sulfonamide; LiN(SO 2 CF 3 ) 2 ) has become quite popular as a replacement salt for LiPF 6 but several issues are still not addressed satisfactory. LiTFSI shows a good thermal stability when compared to LiPF 6 under similar conditions and is not prone to HF formation [239, 240]. LiTFSI solutions show reasonable ionic conductivity but are less conductive than LiPF 6 solutions [58]. The biggest drawback though is the Al corrosion under charged conditions [241]. Recently, it was demonstrated that concentrated solutions of LiTFSI in EC or EC/DEC can suppress the aluminum corrosion but the conductivity in such systems is rather low with 0.1 and 0.5 ms cm 1 respectively [242]. Furthermore, using large amounts of expensive salts seems counter intuitive from an economical point of view. Another approach studied to compensate for the aluminum corrosion was by using small amounts of LiPF 6 in combination with LiTFSI in an EC/DMC solution. Even 1% of LiPF 6 was enough to suppress the Al corrosion [243]. Additives Probably the easiest way to influence the properties of electrolyte solutions can be found by adding small percentages of additives to the solution. The aimed improvements are usually related to SEI stability/formation, cathode protection, LiPF 6 stabilization and safety features such as overcharge protection or flame retardation [245]. Reductive additives which have a higher reduction potential and react before the actual SEI formation takes place, can produce a preliminary surface film on the graphite anode. This layer will serve as a cornerstone for the SEI formation, limiting side reactions and effectively lowering the irreversible capacity loss due to the SEI formation. A popular candidate is the polymerizable vinylene carbonate which shows improvement in irreversible capacity loss and cycle stability with graphite anodes [246, 247]. The additives do not necessarily have to be polymerizable and other studied compounds are known to be beneficial as well (e.g. ethylene sulfite [248]). The idea is that they decompose to insoluble products at a higher reduction potential forming a preliminary surface film as well. Even materials that do not decompose within the operating potential may assist in the SEI formation by scavenging radicals and suppressing side reactions or by reacting with the SEI products to form more stable compositions. One of such compounds is the LiBOB salt which is known to form very stable SEI layers and even as little as 1 5 mol % have been reported to be sufficient to enhance the cycle stability [249]. LiPF 6 stability under elevated temperatures is a big problem in LIB and successfully suppressing or at least inhibiting the HF formation can have a tremendous positive impact on long term cycle stability. One way to shift the equilibrium of the PF 5 formation (LiPF 6 LiF + PF 5 ) can be found by simply adding LiF to the electrolyte solution in a real fundamental usage of the law of mass action [250]. Even in the absence of water, PF 5 is known to react directly with the SEI components thus forming LiF, POF 3 and CO 2 in the process. The gas evolution and especially the formation of the very toxic POF 3 are problematic and result in a buildup of pressure in the cell thus leading to increasingly more dangerous conditions in commercial cells. Lowering the acidity of PF 5 is a possible way to reduce the side reactivity and consequently the gas evolution in LIB. Weak complexing agents (e.g. Tris(2,2,2-trifluoroethyl) phosphite TTFP [251]) can coordinate PF 5 leading to a better cycle stability of LIB even under elevated temperatures. LiPF 6 stability is also directly linked to the cathode stability since deterioration occurs due to the presence of HF. Scavenging water or HF are possible routes to enhance cathode stability and a number of components have been positively tested such as butyl amine [252] or N,N'-dicyclohexylcarbodiimide [253]. Electrolyte solutions can be modified in many ways and even changing only one component can have impacts on many attributes of the whole system leading to a series of necessary tests to truly validate a new solvent or salt. Simply showing certain improvements in one aspect but ignoring the impact on other aspects makes for poor science but has become an increasingly common practice. Electrolyte solution research is very demanding and a broad skill set is required to understand various aspects in details. Especially full cell tests are time consuming and short half-cell tests (~100 cycles) are usually favored but allow only a glimpse of the picture [254]. Dahn et al. recently published a series of papers employing high precision voltage/current/time measurement systems, as an extension to common battery test units [ ]. Whenever the coulombic efficiency (CE) is anything but exactly 1, it means that the system is subject to side reactions. Using high precision CE measurements are an elegant way to predict capacity fading after a few cycles. It was possible to estimate the capacity at cycle 750 of LTO/LCO or graphite/lco cells with various additives after an initial cycling test (1000 h; ca. 35 cycles) providing a way to shorten total test times and circumvent long term cycling to some

16 1110 FLORIAN SCHIPPER, DORON AURBACH Table 3. Overview of some materials studied as anodes in lithium ion batteries, properties taken from ref. [34, ] Material Theoretical capacity, ma h g 1 Potential range vs. Li + /Li, V ρ, g cm 3 25 C Volume change Li % LiC ~ ~10% Li 4 Ti 5 O ~ ~0.2% Li 4.4 Si % Li 4.4 Sn % extend [255]. Considering that commercial LIB can have any number of added compounds in a wide percentage range, the methodology developed by Dahn et al. might be a good option to carry out combinatorial studies. Advancements in Anode Materials Research into anode materials is most often aimed at lithium metal alloys such as Li/Si or Li/Sn (Table 3). Graphite as a real intercalation anode with a low intercalation potential close to metallic lithium was something of a lucky find, as was described earlier in the paper. Since the capacity of graphite is usually twice as high as those of most cathode materials, mass balancing in full cells is easily done. Lithium titanate Li 4 Ti 5 O 12 (LTO) is another intercalation material which can be combined with high voltage materials to give lithium ion battery full cells with reasonable voltages in the range V. Graphite is the dominating anode material in commercial LIB and is currently used either from a synthetic route (mesocarbon microbeads, MCMB [258]) or from modified natural flake graphite [259]. The expected high demand in graphite for LIB in light of electric vehicle development puts natural graphite more and more on the scope of companies since it has the potential to be available for roughly half the price than synthetic graphite [260]. Even though having anode materials such as metallic lithium with > 3000 ma h g 1 sounds very impressive, combing such in a useful manner with cathodes only providing around 200 ma h g 1 is not so important, from energy density point of view. High capacity Li anodes are more suited to be the negative material in lithium sulfur or lithium oxygen batteries which will be discussed later. The next part shows some recent developments in metallic lithium, silicon and tin anodes. Lithium Metal Anode A number of researchers are still working to improve the performance of metallic lithium anodes in hope to exploit their high gravimetric and volumetric capacity. Three problems hinder the implementation of metallic lithium: (1) dendrites formation (leading to short circuits), (2) poor capacity retention when practical specific charge (per cm 2 ) is exchange upon cycling and (3) unavoidable continuous reactions with the electrolyte solutions in practical cells, which dry the batteries [264]. Different approaches to achieve limitation or complete suppression of dendrites growth have been studied and can be divided into SEI related approaches (modifying the Li surface chemistry) or mechanical blocking of dendrites growth [265]. A very well-studied electrolyte solution based on LiAsF 6 in 1,3-dioxolane with stabilizing tertiary amine additives is one of very few systems that was actually proven to be suitable for rechargeable Li batteries [266]. And practical Li/LiMnO 2 cells were commercialized [267]. The improved cycling behavior was attributed to the unique and complicated surface chemistry of Li metal in this electrolyte solutions that forms flexible SEI layer, that can accommodate morphological changes during periodic lithium dissolution/deposition [37, 268]. Unfortunately, the electrochemical stability window is limited to ~3.7 V which made the electrolyte solution a good choice for 3 V LiMnO 2 cathodes but is unsuitable for most of the current generation cathodes [266]. Despite the renaissance in the study of Li anodes in recent years, we are skeptical about their relevance to rechargeable batteries based on liquid electrolyte solutions [37]. It is possible however to use Li metal anodes in solid-state batteries, in which the electrolyte systems are based on polymeric matrices (e.g., based on derivatives of polyethylene oxide [269, 270]) or of solid Li ions conductors (e.g., Li 3 PO 4 LiPON [271]). The fields of solid-state and polymeric electrolytes are broad, well developed and should be beyond the scope of this paper [272]. Silicon and Tin Anode A number of elements (Al, Si, Sn, Sb, Ga, Ge etc.) are known to form alloys with lithium in electrochemical processes, when polarized to sufficient low potentials vs. lithium [273]. The most important elements in this respect are silicon and tin, whose alloying behavior with lithium is very similar to each other.

17 A BRIEF REVIEW: PAST, PRESENT AND FUTURE OF LITHIUM ION BATTERIES 1111 Both silicon and tin can be alloyed with 4.4 moles of lithium leading to a substantial volume expansion during lithiation (see Table 3). This poses a multitude of problems when silicon or tin are used as alloying anodes. Li Si or Li Sn compounds cannot withstand the volume expansion and contraction without breaking apart due to high compression and tensile stress respectively. This leads to cracks within the composite anode films and eventually more and more active mass becomes detached form the conductive matrix resulting in an accumulation of dead mass in composite electrode [274]. Furthermore, the cracks thus formed continuously expose fresh active Li-alloy surface to the electrolyte solution leading to ongoing side reactions which consume the electrolyte solution [275]. It was also shown that lithium extraction is not completely possible in an electrochemical fashion and some lithium remains within the alloy anodes [276]. All the above mentioned processes lead to a rapid capacity fading in alloy anodes and a large irreversible capacity loss is observed during the first cycle (Fig. 12) [277]. Since the volume expansion causes the most problems in alloy anodes, ways to accommodate or alleviate the mechanical stresses are commonly used to enhance cycle stability. It is know that properties of nano-sized materials differ from those of micronsized particles, which also applies to tensile and compressive strength [278]. Shifting the particle sizes of alloy anodes from micron-size to submicron particles increases the cycle stability of the material considerably [279, 280]. Nanoparticles in alloy anodes come with all their advantages (e.g. short diffusion pathways) but also disadvantages (e.g. high surface area, fabrication costs). Furthermore, nano-silicon was shown to have a strong tendency to merge to bigger particles offsetting the initial gain by nano-sizing the material [276]. Imbedding nano-particles in an appropriate host matrices, such as carbonaceous composites, is an easy way to isolate the particles and ensure intimate contact with to the conductive backbone [281]. Excellent cycling stability was reported for amorphous thin film silicon anodes which are produced via magnetron sputtering or vapor deposition [282, 283]. The performance of such thin film electrodes is strongly linked to their thickness and decrease with increasing thickness. This is likely due to longer lithium diffusion pathways and stronger impact of compression and tensile stress in case of too thick alloy films [284]. It is not completely understood why amorphous silicon can withstand the high volume change in a better way compared to their crystalline counterpart. Explanations are often given by the one phase nature found during cycling of amorphous silicon. In contrast, crystalline silicon undergoes a phase change from crystalline to amorphous nature during repeated lithiation/delithiation [276, 285]. Voltage, V vs. Li/Li Discharging (de-alloying) Charging (alloying) Specific capacity, ma h g 1 Fig. 12. Typical voltage profiles for silicon anodes with 5 cycles shown; reprinted with permission form ref. [277], copyright 2005 Elsevier. The main focus in alloy anode research lies in understanding capacity fading mechanisms and to tailor materials accordingly. The high theoretical capacities of Si and Sn are very tempting but not necessarily useful in full cells. Sacrificing capacity in order to gain cycle stability is definitely an option. A stable silicon anode delivering 800 ma h g 1 is much easier to balance in lithium ion batteries than anodes having initially 3000 ma h g 1, but showing capacity fading upon cycling. Higher capacities might still prove useful in lithium sulfur batteries were usually metallic lithium is used as anode in fundamental research but might not be applicable in future commercial cells [286]. The above discussed approaches to achieve better cycle stability are a selected number of strategies found in the literature, but are far from providing ultimate solutions [55]. BEYOND LITHIUM ION BATTERIES 4000 This final chapter will summarize recent efforts to develop very high energy density Li batteries, by giving up Li ions intercalation cathodes and moving to more energetic cathode reactions. We cover herein very briefly the emerging fields: Li-sulfur and Li-oxygen batteries. Low molecular weight cathode materials are very attractive in terms of high specific capacity. Sulfur (M(S): 32 g mol 1 ) and molecular oxygen (M(O 2 ): 32 g mol 1 ) are being extensively studied as cathode materials in combination with metallic lithium anodes. Both systems have been investigated for decades with decisive findings in recent years putting them in the focus of many research groups and companies around the world. The Li S system is more advanced and is believed to reach commercialization within the next 5 10 year while Li oxygen systems are often ball-parked in a year range. 1

18 1112 FLORIAN SCHIPPER, DORON AURBACH Voltage, V vs. Li/Li I II Charge Discharge III S 8 Li 2 S 8 Li 2 S 4 Li 2 S 2 Li 2 S 0% 12.5% 25% 50% Depth of discharge Fig. 13. Typical discharge profile of a lithium sulfur battery (sulfur is loaded onto CMK-3); reprinted with permission from ref. [288], copyright 2010 the Royal Society of Chemistry. Lithium Sulfur Battery (LSB) Publications on LSB can be found dating back to the early 1970 s a time when LIB were still two decades away from commercialization and research on Li battery systems begun to flourish [287]. After LIB were introduced to the market, the work on Li S became less attractive and interesting. The efforts to promote the field of electro-mobility, due to increasing environmental constraints related to the use of fossil fuels renew work on Li sulfur systems, due to their potential high energy density. Innovative work on novel sulfur cathodes was published in 2009 by Nazar et al. [288]. LSB have a multitude of problems which made these systems not very appealing. However, intensive research efforts on these systems seem to provide some solutions to old problems, related to Li S cells [289]. The reaction between lithium and sulfur can be simplified to: 2Li + S Li 2 S which gives a theoretical capacity of 1675 ma h g 1 based on sulfur. This high capacity makes Li S very interesting as a future power source and practical values of 1000 ma h g 1 (S) have been reported [290]. It is important to mention that capacities for LSB are often reported in respect to their sulfur mass and not to the total mass of the sulfur/carbon composites. Also, the sulfur loading in composite electrodes is often in the range of wt % which can be very misleading in terms of actual capacity. The working voltage of a LSB is about 2 V which leads to a theoretical gravimetric energy density of 2300 W h kg 1 (based on a sulfur cathode and lithium anode). Projections often put packed LSB around 600 W h kg 1 a value about three times higher than that of state-ofthe-art LIB. Furthermore, sulfur is the 10th most abundant element on earth and can be easily exploited in near 99.9% purity via the Frasch process [291]. In addition sulfur is a side product of the petro chemical industry and huge quantities are produced in excess every year [292]. LSB do not only have a high possible energy density but the low price of sulfur put them also in better economical standings compared to LIB. Sulfur is non-toxic and environmentally benign. However, problems linked to the sulfur cathode are the low electrical conductivity of sulfur, the volume expansion of sulfur reduction products (LiS n moieties, down to Li 2 S) during charge/discharge and the dissolution of polysulfides into the electrolyte solution during cycling. The latter may undergo continuous reduction processes at the negative electrode, which leads to a shuttle mechanism that avoids completed re-oxidation of the Li-sulfide moieties to sulfur. Due to which the capacity of the sulfur cathode in Li S cells is limited. Sulfur shows an electrical conductivity of ca S cm 1 what makes it a very poor electrode material [293]. Early work on LSB focused either on high temperature approaches or using electrolyte solutions containing lithium polysulfides to overcome the low conductivity of pure sulfur. Thermal lithium sulfur batteries have to operate at >400 C in order to ensure solubility of the short chain polysulfides and sufficient mass transport within the cell. However, such cells are very sensitive to contamination [294]. Better results were obtained for sulfur containing compounds as active material such as FeS 2 and FeS [295]. However, FeS x cathodes require high operating temperatures and suffer from internal resistance. Hence, their importance as practical battery materials is limited. Room temperature cells were tested in the late 1980s by Peled et al. with polysulfides dissolved in THF [296]. The cells were plagued by passivation of the lithium anode with insoluble discharge products (e.g. Li 2 S 2, Li 2 S) which leads to a poor electrochemical performance. The discharge process in LSB progresses from the cyclo-octasuflur (S 8 ) involving polysulfides (Li 2 S n ; n = 2 8) to Li 2 S. The density changes from 2.07 g cm 3 (S 8 ) to 1.64 g cm 3 (Li 2 S) leading to a volume expansion of ca. 80% and eventually to a pulverization of Li 2 S [297, 298]. The contact loss between sulfur and the current collector leads to an increase in dead sulfur in the cells and eventually capacity fading. The volume expansion cannot be circumvented and has to be addressed in case of practical LSB for safety and durability reasons. Furthermore, the Li polysulfides (Li 2 S n, n > 4) are known to be readily dissolved in commonly used electrolyte solutions [296]. A typical discharge profile for LSB displays a stepwise voltage behavior which reflects the gradual reduction of sulfur to long chains polysulfide moieties and then, their reduction to Li 2 S (Fig. 13). Dissolved Li polysulfides migrate through the cell and react with the lithium anode directly to insoluble Li 2 S and Li 2 S 2. This process is in fact a shuttle mechanism that avoids a full charging of Li 2 S back to sulfur, which considerably

19 A BRIEF REVIEW: PAST, PRESENT AND FUTURE OF LITHIUM ION BATTERIES 1113 (а) 1400 (b) 6.5 nm 3 nm Specific capacity, ma h g Cycle number 20 Fig. 14. (a) Schematic representation of sulfur loaded CMK-3; reprinted with permission from ref. [290], copyright 2009 Nature; (b) Cycling performance of CMK-3/S (red points) and PEG modified CMK-3/S (black points); reprinted with permission from ref. [290], copyright 2009 Nature. limits the capacity of Li sulfur cells. Although it was shown that the presence of Li 2 S and Li 2 S 2 in surface films formed on lithium anodes in Li S cells may improve their passivation properties, avoiding contact between sulfur reduction products and Li anodes in Li S cells is very desirable for their successful operation [299]. An elegant way to alleviate problems arising from the flow of LiS n species towards the anodes in Li S cells is found by encapsulating sulfur in carbonaceous matrices, as was suggested by Nazar et al. [290]. The employed carbon CMK-3 [300] (pore volume: 2.1 cm 3 g 1 ) was mixed with sulfur and was heat treated to achieve a melt infusion of sulfur into the meso-porous network. Followed by a final heating to 155 C, in order to evaporate excess sulfur (Figs. 14a, 14b). The intimate contact between sulfur and the electrically conductive carbon framework, compensates for the low conductivity of sulfur. The high pore volume can account for the volume expansion during sulfur reduction. Furthermore, the Li sulfides dissolution is retarded when cathodes comprising sulfur encapsulated in carbon are used for Li S cells. The pioneering work of Nazar et al. promoted extensive work on composite sulfur-carbon composite cathodes. It was shown that dissolution of LiS n species can be suppressed by using functional carbons (e.g. nitrogen [301], sulfur [302]) instead of pure carbon. The pore structure of carbons for sulfur cathodes have to be tailored to meet the needs as viable host networks. Macro-, meso- and micropores are defined to be > 50 nm, 2 5 nm and < 2 nm respectively. Macroporous carbon materials are attractive since they can hold a decent amount of sulfur leading to high specific capacity of Li S cells. Anyway, the performance of all Li S cells which use common electrolyte solutions is poor due to the above described shuttle mechanism. Recently, better cycling performances of Li S cells have been reported for cathodes based on macroporous carbons, by using a room temperature ionic liquids in which the Li polysulfides shuttle is suppressed due to the high viscosity of these electrolyte solutions [303]. The majority of work on carbon hosts structure is devoted to mesoporous carbons with high pore volumes with which high sulfur loadings, of up to 70%, can be achieved [290, 301, 302, ]. The small channels of these carbons lead to a better contact between the insulating sulfur and the conductive carbon matrices. Nevertheless, there is no way to avoid the detrimental shuttle mechanism in Li S cells when any open structure S C composites are used. Several approached to fully mitigate and avoid the shuttle mechanism in Li S cells were suggested and tested in recent years. We mention below several important ones. Polymer coatings on S-carbons matrices have been investigated, with some level of success [307, 308]. Moving from mesoporous carbons to microporous carbons with very small pores (0.5 nm) gives an opportunity to confine short chained sulfur allotropes (S 2 S 4 ) into the microporous structure of the carbon [309]. Nazar s group promotes development of composite sulfur electrodes based on sticky surfaces that strongly adsorb to them the LiS n formed upon sulfur reduction. Several families of suitable substrates were demonstrated for effective composite sulfur cathodes by this highly innovative group [310, 311]. Another approach, promoted by Aurbach s group is to work with S C matrices covered by stable surface films which behave like SEI [312, 313], so sulfur reduction occurs as quasi-solid-state reactions, in which the products remain fully trapped in the porous structure.