Highly Conducting 3D-Hybrid Polymer Electrolytes for Lithium

Article pubs.acs.org/cm

Highly Conducting 3D-Hybrid Polymer Electrolytes for Lithium Batteries Based on Siloxane Networks and Cross-Linked Organic Polar Interphases Nicola Boaretto,†,‡,§ Andreas Bittner,‡,§ Christine Brinkmann,§ Birke-Elisabeth Olsowski,§ Jochen Schulz,§ Mona Seyfried,§ Keti Vezzù,∥ Michael Popall,§ and Vito Di Noto*,† †

Dipartimento di Scienze Chimiche, Università di Padova, via Marzolo 1, I-35131, Padova, Italy Lehrstuhl für Chemische Technologie der Materialsynthese, Universität Würzburg, Röntgenring 11, D-97070 Würzburg, Germany § Fraunhofer-Institut für Silicatforschung ISC, Neunerplatz 2, D-97082 Würzburg, Germany ∥ Veneto Nantotech S.C.p.a., via San Crispino 106, I-35129, Padova, Italy ‡

S Supporting Information *

ABSTRACT: The development of polymer electrolytes with high ionic conductivity, high lithium transference number, and high electrochemical stability is one of the main aims in the field of lithium battery research. In this work, we describe the synthesis and the characterization of new electrolyte systems, composed of three-dimensional hybrid inorganic−organic networks doped with LiClO4. The preparation route comprises only three steps, namely a sol−gel reaction, salt dissolution, and an epoxide polymerization reaction. The lithium concentration, and thus the lithium transference number, was modulated by adding lithium hydroxide in the sol−gel step. In this way, seven electrolytes with varying salt concentrations were prepared. The hybrid electrolytes are characterized by good ionic conductivities (up to 8·10−5 S/cm at room temperature) and high thermo-mechanical and electrochemical stabilities. Stability tests versus lithium metal via galvanostatic polarization showed that this material is superior with respect to reference poly(ethylene oxide) based electrolytes.

W

addressed, either by varying the lithium salt8 or by developing single-ion conduction electrolytes.2,9−11 In the field of the binary polyether based electrolytes, one of the easiest and most rewarding research strategies consists of the preparation of composites, in which inorganic particles are dispersed in the PEO matrix. These systems usually benefit from improved mechanical properties and from an increased ionic conductivity.12,13 The latter effect is attributed to the hindering of the crystallization process or to Lewis acid−base type interactions between the filler surface functional groups, ions, and ether oxygen groups.14−16 In general, because of their large specific area, nanosized particles have greater effects compared to micrometer-sized particles.17 In the 1990s, a new class of hybrid electrolytes was developed, based on sol−gel class II hybrids, ORMOCERs as Polymer Electrolytes (ORMOCERs - APE).18−22 These ion conducting materials consisted of inorganic and organic nanodomains, formed via a sol−gel reaction of functionalized alkoxysilanes. By this method, a highly homogeneous media was obtained, and the

ith the demand for higher energy densities, coupled with the need of increased safety, the electrolyte is considered to be the key component for the development of improved lithium batteries.1 In particular, much effort is currently spent on the development of solid polymer electrolytes (SPEs) which provide higher thermal stability with respect to standard liquid electrolytes. They also offer better resistance to dendrite formation, thus paving the way for the use of lithium metal anodes and to high energy density batteries such as Li/air and Li/S batteries.2 The most studied class of polymer electrolytes consists of complexes of poly(ethylene oxide) (PEO) with various lithium salts, as described by Wright and Armand already in the 1970s.3 These materials are of great interest because of their low cost and toxicity, but the conductivity at room temperature is restricted to ca. 10−6 S/cm only, which is too low for practical purposes.4 One of the causes of the low conductivity is the semicrystalline morphology of PEO: ionic conduction occurs predominantly in the amorphous domains, with the crystalline domains playing an impeding role by increasing the tortuosity of the conduction pathways.5,6 Other drawbacks are the decrease of the mechanical and electrochemical stability at high temperatures2 and the low Li+ transport number.7 This point is particularly critical for electrolyte application, and it has been variously © XXXX American Chemical Society

Received: July 7, 2014 Revised: November 3, 2014

A

dx.doi.org/10.1021/cm5024647 | Chem. Mater. XXXX, XXX, XXX−XXX

Chemistry of Materials

Article

atmosphere. The progress of the reaction was monitored by FT-IR, H-, 13C-, and 29Si NMR spectroscopy. Electrolyte Preparation. P was mixed with different amounts of LiClO4, and hybrid-polymer electrolytes were prepared by epoxy-ring opening reaction, catalyzed by BF3·NH2Et. From now on, the electrolytes will be named as [HP/(LiClO4)x], where x is the molar ratio n(LiClO4)/n(EO), calculated on a basis of 17 mol(EO)·kg−1 for the hybrid matrix. This value was determined by stoichiometric calculations. In total seven sets of electrolytes were prepared, with x equal to 0, 0.005, 0.01, 0.02, 0.05, 0.076, and 0.1 (Table 1).

interaction between organic and inorganic domains was maximized. As a result, the materials showed low crystallinity, high thermomechanical stability, and improved conductivity compared to standard PEO-LiX complexes. The synthesis and properties of some of these hybrid electrolytes were described in previous papers.20−22 In this report, we present a modified synthetic pathway (Scheme 1): first, a liquid precursor is

1

Scheme 1. Three-Step Synthesis of the Hybrid Electrolytes

Table 1. Electrolytes Composition and LiClO4 and Li+ Concentration n(EO)/kg

n(LiClO4)/n(EO)

c(LiClO4)/mol·kg‑1

c(Li+)/mol·kg‑1

0 0.018 0.036 0.072 0.181 0.275 0.362

17 16.85 16.70 16.41 15.59 14.95 14.39

0 0.005 0.01 0.02 0.05 0.076 0.1

0 0.084 0.167 0.328 0.778 1.136 1.439

0.43 0.51 0.59 0.74 1.17 1.52 1.81

To obtain the hybrid-polymer electrolytes, 2% by weight of BF3· NEtH2 was added to the mixture P/LiClO4. The resulting mixture was poured into 18 mm wide aluminum pans, and the samples were heated at 100 °C for 70 h on a heating plate, under Argon atmosphere. After cooling, the free-standing pellets were peeled off and characterized. Instruments and Methods. Thermogravimetric analyses were carried out with a STA 449C Jupiter (Netzsch) thermobalance. The TG profiles were recorded under N2 flow (20 mL/min), in a temperature range from 20 to 950 °C. The temperature ramp rate was set to 10 °C/min. FT-IR ATR spectra in the medium infrared region (MIR) were collected using a Nicolet iS5 spectrometer equipped with a Specac Quest single-reflection ATR accessory (diamond crystal). Each spectrum was recovered by averaging 1000 scans. Baseline correction was performed with the Nicolet FT-IR Nexus spectrometer software. Raman spectra were collected with a Bruker RFS 100 FT-Raman spectrometer. The spectrometer is equipped with an air cooled 500 mW Nd:YAG-laser with wavelength of 1064 cm−1 and a liquid nitrogen cooled Ge detector. The measurements were obtained by averaging 1000 scans. NMR spectra were obtained with a Bruker AMX 400 and with a Bruker DSX 400 spectrometer (400 MHz). Rheological analysis of [HP/(LiClO4)0] was performed with an Anton Paar MCR 502, equipped with a PP-25 parallel plate measuring system. In a typical experiment, a pellet of diameter 25 mm was subjected to a sinusoidal angular deformation of amplitude 0.18 mrad, with frequency varying from 0.1 to 10 Hz. The sampling frequency was of 10 points/dec, and the normal applied force was set to 5 N. The experiment was performed in the temperature range between −20 and 100 °C, at an interval of 10 °C. Gel permeation chromatography (GPC) was performed with a Gamma Analysen Technik TC 1900 equipped with PSS column and refractive index detector. THF was used as eluent, and the measurement temperature was set to 38 °C. ICP-AES measurements were performed with a Varian WISTA PRO spectrometer, with a lower detection limit of 0.02 mg/L. Linear voltammetry experiments, transport number measurements, and lithium stability tests were performed with Commercial EL-ECC cells and with a Solartron Analytics 1470 E potentiostat, equipped with a frequency response analyzer (FRA). Impedance measurements were performed in the frequency range from 1 MHz to 100 mHz, with an AC amplitude of 20 mV. Linear sweep voltammograms were collected using gold as working electrode and lithium as counter and reference electrode. The high and low voltage sweeps were collected separately at a scan rate of 0.5 mV/s on different samples, in the voltage range from 3 to 5 V and from 3 to −1 V (vs Li+/Li), respectively. The lithium transport numbers were determined by the Vincent−Bruce

obtained by hydrolysis/co-condensation reaction of a polyethylene glycol-α-methyl, ω-propyltrimethoxysilane ether (Me(PEG)8-PTMS) together with a cross-linker, namely 3(glycidoxypropyl) trimethoxysilane (GPTMS). With respect to the previously reported approach, the new synthetic pathway has the advantage of greater simplicity, as the precursor is obtained in only one step. Furthermore, no chlorinated compound is used during the synthesis, which, if not completely removed, may adversely affect the cell performance. The solid electrolyte is obtained by mixing the precursor with a suitable lithium salt (LiClO4 in this case) and by cross-linking via the epoxy-moieties of the GPTMS. The resulting structure is composed of cross-linked oligo-siloxane networks with pendant polyether chains. The high degree of mixing between inorganic and organic domains inhibits the reorganization (i.e., crystallization) of the polyether domains. In this way, a good compromise between thermo-mechanical stability and ionic conductivity is achieved.

■

LiClO4/g

EXPERIMENTAL SECTION

Materials. PEG (8) -α-methyl,ω-propyltrimethoxysilane (Me(PEG) 8 -PTMS, Gelest, 98%) and 3-(glycidoxy-propyl)trimethoxysilane (GPTMS, Sigma-Aldrich 98%) were used as received. LiOH (Sigma-Aldrich, 98%) and LiClO4 (Sigma-Aldrich, 99.99%) were dried for 24 h at 150 and 120 °C, respectively. All other reagents and solvents were used after purification by standard methods. All reagents were stored under argon over molecular sieves to prevent contamination by moisture. All transfer and handling operations were performed under Argon atmosphere. Precursor Synthesis. In a typical synthesis, 15.21 g of Me(PEG)8-PTMS (0.029 mol), 0.56 g of distilled water (0.031 mol), and 0.26 g of dried LiOH (0.011 mol) were dissolved in 10 mL of dry diethylcarbonate (batch A). The mixture was stirred overnight at room temperature. In another flask (batch B), 2.36 g of GPTMS (0.010 mol), 5 mL of diethylcarbonate, and 0.27 g of distilled water (0.015 mol) were mixed and stirred overnight at room temperature. Batch B was added to batch A, and the mixture was stirred at room temperature for another 24 h. The solvent was evaporated, and the product (P), a slightly colored viscous liquid, was dried under high vacuum, at 50 °C for another 24 h. P was finally stored under argon B



Article

method.23 A symmetric cell with lithium as working and counter electrode was assembled and stored at 50 °C for 24 h. After cooling to room temperature, an overpotential of 20 mV was applied to the working electrode. The resulting current flow was monitored, and the polarization was maintained until a steady state was reached. In general, more than 100 h were necessary to reach the steady state. Impedance spectra were collected both at the beginning and at the end of the chronoamperometric experiment. The lithium transport number (t+) results from the equation

t+ =

Is (ΔV − R 0I0) · I0 (ΔV − R sIs)

silanes and were thus expected to hydrolyze easily even without catalyst. However, previous trials proved that K1 has a lower hydrolysis rate, probably because of the bulky pendant PEG chain and the interaction between the oxyethylene and polar SiOX functions. To avoid self-condensation of K2, K1 was therefore allowed to prehydrolyze for 24 h. The cocondensation of the two components was expected to result in the formation of three-dimensional oligosiloxane cages with pendant polyether fragments. These units could be later crosslinked through the epoxy moieties, therefore resulting in a highly disordered three-dimensional structure composed by intimately mixed polyether and siloxane domains. During the precursor synthesis, LiOH was also added. This was done in order to achieve two goals: a) the catalysis of the cocondensation reaction and b) the formation of lithium silicate bonds. The latter point was supposed to lead to an increase of the lithium concentration with respect to the mobile anions and therefore to a possible increase of the lithium transference number. Furthermore, the formation of lithium silicate bonds would have lowered the overall siloxane condensation degree, thus increasing the structural disorder and possibly increasing the overall conductivity. To achieve these objectives and to hinder the formation of electrochemically unstable silanols, the water content was kept as low as possible, in a ratio 1:2 to the alkoxy-silicon bonds. As it will be shown later, the result was rather unexpected, since the presence of lithium hydroxide, the choice of diethylcarbonate as the reaction solvent, and the low concentration of water lead to a predominance of a parasite reaction, e.g. a transesterification reaction with the formation of silicon-ethoxy bonds. P and K1 were characterized by 1H-, 13C-, and 29Si NMR spectroscopy. The spectra of GPTMS are not shown, but the complete assignation can be found elsewhere27 and it is reported in Table 2. Starting with the 1H NMR (Figure 1), the spectrum P shows the peaks originating from the alkyl/ polyether protons (0.60, 1.68, 3.37, 3.43, 3.65 ppm) and the peaks of the glycidoxy-propyl fragment (1.56, 2.59, 2.78, and 3.13 ppm). The peak characteristic of the K1 silicon-methoxy protons is detected at 3.55 ppm, and its intensity is strongly decreased in P. This could indicate the presence of residual silicon-methoxy groups, but this possibility is ruled out by the absence of the relating peak in the 13C spectrum (Figure 2). The contradiction is resolved by 1H/13C COSY (Figure S1) which shows that the methoxy proton signal (3.55/50 ppm) overlaps the signal of a proton belonging to the PEG chain (3.356/71.9 ppm). The absence of the methoxy couple in the P spectrum confirms that the silicon-methoxy groups have reacted completely. In the 1H spectrum of P, two new peaks appear, at 1.24 and 3.80 ppm. These peaks correspond to a silicon-ethoxy fragment, respectively, in positions β and α to the oxygen. The analysis of the peak intensities indicates a relative population α/β of about 2/3. Compared to the peak centered at 0.6 ppm (proton in position α to the silicon), the intensities of the peaks at 1.24 and 3.80 ppm are doubled. This shows that there are at least two bonds Si−O−C for every Si−C bond, i.e. for every silicon atom. As anticipated, the origin of these groups is probably an alcoholate exchange reaction with the solvent (diethylcarbonate), catalyzed by LiOH (Scheme 2). The other peak intensities agree with the reaction stoichiometry. In particular, the spectrum indicates a ratio of 0.25 epoxy-protons per Si-CH2 group, again in agreement with the stoichiometry of

(1)

where Is and I0 are the steady state and initial current, respectively, ΔV is the applied polarization, R0 is the initial interface resistance, and Rs is the steady state interface resistance. The temperature control was guaranteed by placing the cell in a climate-chamber (Weiss, WK3) at a temperature of 25 ± 0.5 °C. The lithium stability test was performed by galvanostatic polarization in a Li/Li symmetric cell.24 A sample was sandwiched between two lithium metal electrodes, and the cell was placed in the climatechamber at 25 °C. A constant current density of +0.004 mA/cm2 was applied for 1 h, then the polarization of the electrodes was reversed, and the same current density was applied for another hour. The loop was repeated 200 times, and, at the end of each cycle, an impedance spectrum was collected. The aim of the experiment was to monitor the development of the cell resistance under dynamic conditions. The complex conductivity spectra were recorded in the frequency range between 10 mHz−10 MHz, using a Novocontrol Alpha-A analyzer. The temperature range from −100 to 100 °C was explored by using a homemade cryostat operating with a N2 gas jet heating and cooling system. The measurements were performed using a sealed homemade cell equipped with platinum electrodes. The geometrical constant of the cell was determined by measuring the electrode−electrolyte contact surface and the distance between the electrodes with a micrometer. The temperature was measured with accuracy better than ±0.02 °C. The bulk conductivity of the materials, σDC, was determined by measuring the conductivity value interpolated in the medium frequency plateau of the σ′(ω) profiles, as described elsewhere.25,26 The performance of a Li/[HP/(LiClO4)0.05]/LFP cell was evaluated by cycling at 60 °C and 0.1 C (calculated on the theoretical specific capacity of LFP). The cathode/electrolyte assembly was prepared by dropping the salt/precursor mixture on a homemade, 18 mm diameter, LFP electrode, and by curing in situ, as described previously. A 16 mm lithium disk was then applied directly on the surface of the electrolyte, and the assembly was mounted in a coin cell. The LFP electrode had a nominal specific charge of 0.0025 g/cm2, corresponding to a nominal specific capacity of 0.425 mAh/cm2. The cell was cycled at constant current (0.1 C, 0.0425 mA/cm2) between 3.0 and 4.0 V. Complete charge was ensured by introducing a constant voltage step at the end of each cycle. This step was set to terminate as the current reached the value of 0.005 C.

■

RESULTS AND DISCUSSION

Synthesis and Characterization of the Precursor (P). The first step of the synthesis was the sol−gel reaction between a Me-PEG functionalized trimethoxysilane, K1, and a crosslinking agent, K2 (Chart 1). Both reagents were methoxyChart 1. a) Me-(PEG)8-PTMS (K1) and b) GPTMS (K2)

C



Article

Table 2. 1H and 13C Chemical Shifts of K1 and K2, with the Assignments Referred to Chart 1 group position K1 K2

δ δ δ δ

1

H C 1 H 13 C 13

1

2

3

4

5

6

7

0.66 5.18 0.58 5.45

1.69 22.68 1.59 23.16

3.43 73.31 3.4 73.61

3.55−3.65 70−72 3.23; 3.64 71.88

3.38 58.94 3.07 50.91

3.56 50.42 2.67; 2.48 44.02

/ / 3.45 50.41

Figure 2. 13C NMR spectra of K1 (a) and P (b). The assignment is referred to Table 2. The peaks α and β in the spectrum P belong to the carbons of a Si−O−Et group (α and β with respect to the oxygen). The peaks “ep” in the spectrum P were assigned to the two carbons of an epoxide ring.

Figure 1. 1H NMR of K1 (a) and P (b). The assignment is referred to Table 2. The peaks α and β (b) were assigned to the protons of a Si− O−Et group (α and β with respect to the oxygen). The inset shows the peaks of the three epoxy protons in the spectrum of P.

the reaction. We can therefore at least exclude the occurrence of any extended ring opening reaction. The 13C NMR (Figure 2) confirms the suggestions obtained by the 1H spectra. Peaks corresponding to the epoxy groups were found at 44 and 51 ppm (carbons 6 and 5 in Chart 1), while the silicon-methoxy characteristic peak at 50.4 ppm was not detected. The relatively intense peaks at 18.3 and 58.3 ppm correspond to the β- and α-carbon of the silicon-ethoxy groups, which have been found also in the proton spectra. The 13C signals, like the 1H signals, exclude the occurrence of any ring opening reaction, since this would result in the formation of a variety of species whose signatures were not detected in the spectrum.28 The 29Si NMR (Figure 3a) reveals four peaks at δ −44, −45, −53, and −61 ppm. The first two were assigned to T0 species (T0 and T0′), arising from noncondensed K2 and K1. The other peaks belong to T1 and T2 siloxane units. By integrating the three signals we calculated the molar fraction of each species, and from this value it was possible to determine the condensation degree, which is the fraction of Si−O−Si bonds over the total possible Si−O−X bonds. The condensation degree for P is equal to 0.31, which means roughly two

Scheme 2. Transesterification Reaction between Methoxysilanes and Diethylcarbonate

noncondensed bonds per silicon. It is also noteworthy that this value corresponds to the concentration of silicon-ethoxy bonds calculated from the 1H NMR spectrum. From these results, we conclude that the noncondensed siloxane bonds are mainly silicon-ethoxy bonds and partly lithium silicate bonds. The molar mass distribution of P was analyzed through GPC (gel permeation chromatography). The elution profile (Figure S2) shows four peaks centered at 400, 1200, 2000, and 4000 g/ mol and a very broad tail going up to 104 g/mol. The first peak was assigned to noncondensed units (nominal average molar mass = 475 g/mol) and the other peaks to oligomeric siloxane units, with the tail indicating the presence of higher mass aggregates with up to 20 siloxane units. The area under the first peak is ca. 10% of the whole spectrum, underestimated compared to what emerged from the 29Si NMR. The mass D



Article

atoms). P contains also 3000 ppm of lithium (0.43 mol/kg) as lithium silicate. Preparation of the Hybrid-Polymer Electrolyte. The electrolytes [HP/(LiClO4)x] were prepared by dissolving different amounts of LiClO4 in P followed by subsequent ring opening polymerization. BF3·NH2Et was used as initiator. By rising the temperature over 90 °C, gas evolution was observed (BF3·NH2Et decomposition), followed by a slow gelation. Most of the viscosity increase was observed within the first 24 h, but reproducible results were obtained only after longer reaction times (up to 70 h). In general, the polymerization rate increased with the salt concentration, due to the increasing polarity of the system. The polymerization process was followed by 29Si NMR, ATR-FTIR, and Raman spectroscopy. A decrease of two characteristic epoxide bands, the ring breathing band at 1254 cm−1 and the CH2 stretching band at 3048 cm−1, was indeed observed in the Raman spectra (Figure 4). Another interesting feature was observed by solid

Figure 3. a) 29Si NMR spectrum of the condensation product P. The relative intensity was normalized over the total area of the four peaks. b) Solid state 29Si NMR spectrum of [HP/(LiClO4)0.05] after the curing step. Figure 4. Decrease of the characteristic epoxide bands after polymerization, observed by Raman spectroscopy.

average molecular mass is equal to 1712 g/mol, whereas the number-average molecular mass is equal to 556 g/mol. The lithium concentration in P, as determined by AES-ICP, was found to be around 3000 ppm (0.43 mol/kg). This value corresponds to the nominal lithium concentration. The position of this lithium in the structure is not easily identified, since there are several possibilities to consider. First of all, the lithium can be present as a salt, either in the form of hydroxide or as carbonate (from solvent decomposition). Second, it can be complexed by the polyether chains (which would require the presence of a counterion); third, it can be present as free alcoholate; and, finally, it can form lithium silicate bonds. Since the IR spectrum (Figure S3) does not show the presence of hydroxides, and the NMR spectrum does not show the carbonate peak, we excluded the presence of both these species. We can also exclude complexation, since the corresponding Li− O stretching band, which would be visible in the Raman spectrum at 865 cm−1 (Figure S4), was not detected.29 The presence of free lithium alcoholate is highly improbable since these species are incompatible with the epoxy groups, which were detected by NMR spectroscopy. For these reasons, the formation of lithium silicate bonds is the only reasonable conclusion for the nature of lithium cations in the precursor. To summarize, the condensation product P is composed of a mixture of noncondensed ethoxy-silanes and oligo-siloxanes of less than 20 units. The condensation degree is about 30% molar, due to the occurrence of a parasite transesterification reaction catalyzed by LiOH. The epoxy groups are still present in the product and their amount, calculated by 1H NMR, is in agreement with the calculated value (1 group every three silicon

state 29Si NMR (Figure 3b): differing from the NMR spectra of the precursor, only one broad peak centered at −67 ppm was observed. This peak corresponds to the absorption of not resolved T2 and T3 species. The absence of signals characteristic of T0 and T1 species, which were observed in the precursor’s spectrum, indicates extended condensation during the curing step. This process is quite surprising since it proceeds in complete absence of water. Presumably, the reaction involves a direct attack of a Si−O−Li+ group on a Si−O−Et group, probably after an attack of BF3 on the alcohoxy-silane oxygen or, alternatively, after a nucleophilic attack of NH2Et on the silicon. An important role is played by the high temperature and, also, by the absence of a solvent. Indeed, the consequent structural organization causes phase segregation between PEGrich domains and siloxane domains, thus bringing the reactive groups in proximity to each other. Vibrational Spectroscopy Analysis. The structure of the seven electrolytes was characterized by FTIR-ATR and Raman spectroscopy. The assignment of the bands is reported in Table 3. Figure 5a shows the Raman spectra of the gel polymer electrolytes in the region between 200 and 1700 cm−1. The features of these spectra can be roughly divided in two groups: the bands belonging to the first group, which are slightly affected by the salt concentration and are due to the vibrational modes of the host matrix and in particular to the Me-PEG fragments; the bands of the second group, which are strongly dependent on the salt concentration and are due to the LiClO4 salt. E



Article

conformation (T = trans, G = gauche) for PEO chain fragments and the formation of helical structures with internal rotation angles of τ(CH2−CH2) = 60° (gauche) and τ(CH2−O) = τ(O−CH2) = 191.5° (trans).30 Major differences are observed in the region of the CH2 stretching, between 2800 and 3000 cm−1 (Figure 5b). In particular, the spectra show a progressive shift to higher frequencies of the νs(A1) (from 2878 to 2885 cm−1) and the growth of a high-frequency side shoulder around 2915−2920 cm−1, with the appearance of a clear splitting for Li:O > 0.05. The shift to higher frequency can be ascribed to the loosening of the C−O bond, which in turn determines a strengthening of the polyether C−H bond. More straightforward is the analysis of the perchlorate peaks, which is carried out on difference spectra obtained by subtracting the spectrum of the sample [HP/(LiClO4)0] from the spectra of the hybrid/salt complexes (Figure 6). Prior to subtraction, all the spectra were normalized with respect to the band at 1470 cm−1, which is insensitive to salt concentration.29 The perchlorate anion, with Td symmetry, has four active Raman modes: the total symmetric stretching vibration ν1(A1), the doubly degenerate mode ν2(E), and two triply degenerate modes, ν3(F2) and ν4(F2).29 Of these four fundamental Raman active modes, just three have been detected: ν1 at 930 cm−1, ν2 at 457 cm−1, and ν4 at 623 cm−1. Furthermore, two other peaks are visible: the peak at 909 cm−1 was assigned to the overtone 2ν2(A1 + E) in Fermi resonance with the ν1(A1) and the peak between 860 and 865 cm−1 which stems from the symmetric Li−O breathing mode. This peak is especially important because it directly provides evidence of the cation complexation by the polyether chains. Indeed, we found no such peak at n(LiClO4)/n(EO) = 0. This means that the lithium which was

Table 3. IR and Raman Vibrational Bands, with the Relative Assignment,29,30 of the [HP/(LiClO4)x] Electrolytes IR wna/cm−1

Raman assign.b

wna/cm−1

Bands Originating from Matrix Vibrations 2940 (sh) νs (CH3) 2920 (vs) 2870 (s) ν (CH2) 2877 (vs) 2840 (sh) νa (CH3) 1470 (m) 1455 (w) sr (CH2) 1456 (m) 1350 (w) w (CH2) 1288 (m) 1300 (w) t (CH2) (E1) 1242 (m) 1253 (w) t (CH2) (A2) 1129 (m) 1200 (w) w (CH2) 1032 (m) 1145 (sh) ν (CO) (E1) 865 (w) 1096 (vs) ν (CO) (A2) 843 (m) 1040 (sh) ν (Si−O) 806 (m) 950 (m) r (CH2) (A2) 850 (m) r (CH2) (E1) 790 (m) ν (Si−O−Si) Bands Originating from ClO4− Vibrations 623 (m) ν2 (ClO4−) 945 (vw) 930 (s) 908 (m) 623 (m)

assign.b νa (CH2) νs (CH2) sr (CH2) }t (CH2) t (CH2) ν (CO) ν (CO) ν (Li−O) }r (CH2)

ν1cip (Li−ClO4−) ν1 (ClO4−) ν3 (ClO4−) ν2 (ClO4−)

a

Wavenumber; relative intensities are reported in parentheses; vs: very strong; s: strong; m: medium; w: weak; vw: very weak; sh: shoulder. b ν, stretching; w, wagging; t, twisting; r, rocking; sr, scissoring.

The position of the PEG bands is in accordance with previous vibrational studies, which indicated a preferential TGT

Figure 5. Raman spectra of the seven hybrid-polymer electrolytes. The molar ratio n(LiClO4)/n(EO) is indicated in the spectra. Each spectrum was normalized to the band at 1470 cm−1. F



Article

Figure 6. Subtracted Raman spectra of the six [HP/(LiClO4)x] (x > 0), in the region between 200 and 1600 cm−1. Each Raman spectrum was normalized on the band at 1470 cm−1 and then subtracted to spectrum for n(LiClO4)/n(EO) = 0. The inset shows an example of peak decomposition of the ClO4− characteristic region between 800 and 1000 cm−1.

introduced during the precursor’s synthesis is not complexated by the polyether chains. This indirectly confirms the formation of lithium silicate bonds. Another interesting feature is the structure of the ν1 absorption around 930 cm−1. This peak is connected to the stretching vibration of free perchlorate anion, while contact ion pairs and larger ion aggregate species are expected to absorb at about 940 and 950 cm−1, respectively.31 In the reported spectra, the peak at 940 cm−1 appears as a weak shoulder in the three most concentrated samples. In order to get more insight into the dissociation state of the salt, the difference spectra in the region between 800 and 1000 cm−1 were decomposed with Gaussians. The inset in Figure 6 shows an example of such decomposition. The figure shows four peaks at 865, 910, 930, and 945 cm−1, assigned to the Li− O breathing, ClO4− 2ν2 mode, to the “free” ClO4− ν1 mode, and to the “paired” ClO4− ν1 mode, respectively. The first three peaks were found in all six difference spectra examined, while the last band occurred only for a molar ratio x ≥ 0.05. In addition, the spectra for x = 0.076 and x = 0.1 showed a broader absorption beyond 940 cm−1, whose decomposition required the use of two other peaks, centered at 955 and 965 cm−1. These two peaks may indicate the presence of dimeric or larger ionic aggregate species, but the low intensity of these bands and the high uncertainty on the baseline prevented us from going too much into detail. For that reason, the group of peaks beyond 940 cm−1 was considered altogether. The areas of the first three peaks showed a monotone increase as a function of salt concentration. In particular, the dependency was roughly linear for the peaks at 865 and 910 cm−1, while the peak at 930 cm−1 showed some deviation from linearity for x ≥ 0.05 (Figure 7a). These deviations are compensated by adding the contribution of the band centered at 945 cm−1. This is an indirect confirmation of the assignment reported and prompted us to consider the ratio of the area of the two bands, ρ = A(945)/ A(930), as an indication of the relative contribution of ion pairing. The ratio ρ, depicted in Figure 7b, is zero up to x = 0.02, while it increases linearly for x ≥ 0.02, reaching a value of 0.17 for x = 0.1. We thus conclude that the perchlorate anion is present mostly as free anion for x lower than 0.05, while, for higher concentrations, the fraction of ionic aggregates increases up to a 17% for x = 0.1.

Figure 7. a) Relationship of A(υ1) and A(υ1+υ1′) with the molar ratio n(LiClO4)/n(EO). b) Relationship between A(υ1′)/A(υ1) = ρ and n(LiClO4)/ n(EO). υ1′ = 945 cm−1 and υ1 = 930 cm−1.

The FTIR spectra (Figure 8) are dominated by the contribution of the polyether chains, especially by the strong C−O stretching bands at 1090 and 1135 cm−1. The position of most of the bands is in agreement with the assignment reported for PEG chains arranged in a TGT configuration.30 The major effect of the LiClO4 addition is the appearance and growth of the perchlorate ν4 band at 623 cm−1. The intensity of this band is directly proportional to the salt concentration (Figure S5a). Another effect is the shift of the asymmetric C−O stretching vibration to lower frequencies, from 1090 to 1070 cm−1 (Figure S5b). This band shift reflects the lithium complexation and the resulting loosening of the C− O bond. The dependence in this case is only linear for low salt concentration (up to x = 0.02), probably reflecting the formation of contact ion pairs, as already observed in the Raman spectra. Finally, in the region of the methylene stretching, we observe a broadening of the symmetric CH2 stretching band, followed by the growth of a band centered at about 2920 cm−1. Again, this effect was already observed in the Raman spectra, and it is connected with the lithium complexation by the polyether chains. In conclusion, the vibrational spectroscopy analysis (Raman and FTIR) reveals several details about the matrix structure, the ion’s dissociation state, and about the interaction between polymer host and ions. Regarding the matrix structure, we observe that a) the PEG fragments assume the typical TGT configuration and form helical structures, b) the configuration is only slightly affected by the salt concentration (at least in the salt concentration range here examined), and c) lithium complexation results in a progressive loosening of the C−O bonds, mirrored by a G



Article

Figure 8. ATR-FTIR spectra of the seven [HP/(LiClO4)x] electrolytes.

strengthening of the polyether C−H bonds. The FTIR spectra show also the bands typical of the Si−O vibrational modes, but a detailed analysis is hindered by the overlap with the much stronger polyether bands. The Raman spectra reveal also some details about the ions’ positions and about the salt dissociation state: a) In the [HP/(LiClO4)0] composition (pure HP), the base lithium is not complexated by the polyether oxygens; therefore, it is probably located in proximity of the silicon−oxygen groups. b) No contact ion pair was observed for n(LiClO4)/n(EO) ≤ 0.02, thus the salt can be regarded as completely dissociated. c) The contact ion pair concentration increases linearly for x ≥ 0.05. This last point, as will be stressed in the next section, has important consequences regarding the ionic conductivity of the electrolytes. Thermal, Rheological, and Electrochemical Characterization. The thermal stability of the seven electrolytes was studied by thermogravimetric analysis in inert atmosphere. The TGA profiles (Figure 9) show no mass loss up to 250 °C. In the case of the [HP/(LiClO4)0], the polyether chains degrade in one step, with onset at 300 °C. On the contrary, for n(LiClO4)/n(EO) > 0, the thermogravimetric profiles show a twostep thermal degradation, at 250 and 300 °C. The strength of the first step increases with the salt concentration and can be explained by polyether chains decomposition due to oxidation reaction with the perchlorate anion. The viscoelastic properties of [HP(LiClO4)0] were studied by oscillatory rheological analysis, in the temperature range between −20 and 100 °C. Figure 10 shows the storage modulus of [HP/(LiClO4)0]. The material’s response can be divided in

Figure 9. TGA profiles of the seven [HP/(LiClO4)x] electrolytes. The thermogravimetric profiles were collected in inert atmosphere (N2).

Figure 10. Storage shear modulus profile of [HP/(LiClO4)0], as a function of frequency from −20 to 100 °C.

H



Article

two temperature regions, below and over 20 °C. In the low temperature region, the elastic shear modulus is higher than 100 MPa. Between 0 and 20 °C, the modulus drops to 50 kPa. Over 20 °C, the elastic shear modulus decreases smoothly with increasing temperature, reaching a value of ca. 20 kPa at 100 °C. The modulus drop just below room temperature corresponds to first order transitions observed in some block copolymers and polymer blends.32 The dependence of G′ on frequency differs also between the two regimes: in the low temperature region, the dependence follows over the whole frequency range the power low G′ ∝ ωα (where ω is the radial frequency), with α comprised between 0.05 and 0.1 (except at −10 °C, where α is higher than 0.2). In the high temperature region G′ follows also a power law dependence of the same kind, but the response can be distinguished between a high frequency region, with α ∼ 0.05, and a low frequency region where the frequency dependence is more marked (α ∼ 0.08). The two regions are divided by a well visible knee which shifts to higher frequencies with increasing temperature. Given the low dependency on frequency, this feature suggests the occurring of a mechanical relaxation. In this sense, the viscoelastic behavior of the material resembles that of a crosslinked polymer, with storage modulus almost independent of frequency. No sign of mechanic degradation up to 100 °C was revealed. The electrochemical stability window of the [HP/(LiClO4)x] complexes was determined by linear sweep voltammetry (Figure 11). All the samples show good stability up to over 4 V (at least 4.5 V for the samples with x ≥ 0.05). At voltages lower than 0 V, the reversible lithium plating branch is visible, whose intensity depends on the electrolyte resistance. In conclusion, all samples show a good electrochemical stability in the range from 0 to 4 V versus Li/Li+. This suggests that these

systems can be used as electrolytes in lithium metal cells, at least in cell configurations where the charge cut off voltage does not exceed much 4 V versus Li/Li+. The stability versus lithium metal was further verified by a galvanostatic polarization experiment. A symmetric Li/Li cell with electrolyte [HP/(LiClO4)0.02] was assembled, and a dynamic stability experiment was performed (see the Experimental Section). The contact stability versus lithium was monitored by recording periodically the impedance spectra of the cell. The impedance spectra in the complex plane resemble a depressed half circle, whose intercepts with the real axis, at high and low frequencies, correspond to the bulk resistance of the electrolyte and to the total resistance of the cell, respectively. An example of this projection is shown in Figure 12a. The difference between the two values gives the

Figure 12. a) First and last impedance spectra (projections in the complex plane) of a symmetric Li/[HP/(LiClO4)0.02]/Li cell during the dynamic lithium stability test. b) Development of the cell resistances (bulk electrolyte resistance Rb and interface resistance Rint) during the same test. The interface resistance of a Li/[PEO/ (LiClO4)0.05]/Li cell, monitored during the same test, is also shown for comparison.

total charge transfer resistance between the two electrodes and the electrolyte, and the development of this parameter was taken as an indicator of the interface stability versus lithium electrodes. The area-normalized values of the interface and bulk resistance are shown in Figure 11b. The bulk resistance remains stable at ca. 5·103 Ω/cm2 during all the experiment, whereas the interface resistance shows an increase from the initial 5.5 × 102 Ω/cm2 to 18·102 Ω/cm2 after 200 cycles. A visual inspection of the lithium electrodes at the end of the experiment showed that the electrodes were still translucent, indicating that only limited lithium oxidation occurred. In this case, the major contribution to the cell resistance is always given by the electrolyte bulk

Figure 11. Linear voltammograms of [HP/(LiClO4)x] in the range between −1 and 5 V. I



Article

describes the contribution of the lithium cations to the charge transport. For a binary salt dissolved in a polymer solvent, the transference number is described by the following equation7

resistance, but in a full cell this value can be lowered by reducing the electrolyte thickness. For comparison, a standard electrolyte complex [PEO/ (LiClO4)0.05] was analyzed as well (Figure 12b). In this case, the interface resistance shows a 10-fold increase, reaching a plateau value of over 7·104 Ω/cm2 after 80 cycles. The increased stability of the [HP/(LiClO4)x] systems, with respect to the PEO based system, may relate to the absence of hydroxyls, which are supposed to behave as preferential sites for attack by lithium metal. At the same time, similar experiments performed on PEG-dimethyl ether based electrolytes showed a behavior similar to PEO based systems, thus indicating that PEG fragments react with lithium metal also in the absence of hydroxyl groups. The observed electrochemical stabilization might also be regarded as a kinetic effect, originating from the restricted diffusion of the reactive species in a reticulated 3Dnetwork. This hypothesis is also suggested by the lack of a clear plateau in the resistance profile of [HP/(LiClO4)0.02]. To summarize, the synthesized hybrid electrolytes show much higher interface stability versus lithium metal compared to classic polyether based electrolytes and are thus potential candidates for application in lithium metal batteries. The ionic conductivity of the seven samples was determined by broadband electric spectroscopy (BES). Figure 13 shows the

t+ = 1 − t− = +

−

z+D+ z D + z −D− + +

(2) +

−

where t and t are the transference numbers, z and z are the charges, and D+ and D− are the diffusion coefficients of cation and anion, respectively. The lithium transference number varies from 0 to 1 and generally, for a polymer electrolyte, is smaller than 0.5.7 As an example, for PEO/LiTFSI complexes, the lithium transference number is close to 0.2.8 Having an electrolyte with a transference number close to 1 is of fundamental importance for the performances of lithium batteries, since small t+ result in the buildup of anion concentration gradients which in turn limit the current densities that can be used in the cells. Indeed, high current densities would cause larger concentration gradients and, in the end, the risks of salt depletion at the cathode and precipitation at the anode.33 Furthermore, low transference numbers are suspected to be responsible for the lithium dendrites growth in lithium metal cells,2,34 a major obstacle to the application of this class of devices. In this study, the Li+ transference numbers at 25 °C were determined by the Vincent−Bruce method23 and are reported in Figure 14. In comparison to the standard PEO/LiClO4

Figure 13. Conductivity profiles versus inverse temperature of the seven [HP/(LiClO4)x] electrolytes.

logarithm of the ionic conductivity (logσDC) for the seven samples, in the range from −100 to 100 °C. The dependence of the ionic conductivity on temperature is complex. Below −70 °C, the ionic conductivity is practically absent. Between −70 and 20 °C, the conductivity increases with an Arrhenius dependence for x ≤ 0.02 and with a VTF (Vogel−Tamman− Fulcher) dependence for x > 0.02. Above room temperature, all samples show a VTF behavior. The discontinuities may relate to thermal transitions (glass transition or melting of crystalline domains). The detailed analysis of the electric spectra and of the relationship between thermal and electrical properties will be presented in another publication. Regarding the conductivity values, they increase with salt concentration and reach a maximum for x = 0.05. At higher salt concentrations, the conductivity decreases strongly, probably due to ion pairing effects and saturation of the hopping sites. The values of conductivity at representative temperatures (−20, 30, and 100 °C) are reported in Table S1. A peculiarity of binary polymer electrolytes is that transport properties are not fully described only by the bulk conductivity. To have a complete picture, one has to determine also the lithium transference number as well, a parameter which

Figure 14. Relationship between the lithium transport number at 25 °C (t+) and the molar ratio cClO4‑/cLi+. The transport number was determined by the Vincent−Bruce method.23 The dotted line was added as a guide for the eye.

complexes, higher transference numbers are expected, because of the higher concentration of lithium cations with respect to the perchlorate anions. Due to the high electrolyte resistance, the value for x = 0 could not be measured. However, since the only mobile charge carrier present is the Li+, a nominal value of 1 is supposed. At x = 0.005, with a lithium concentration six times higher than the perchlorate concentration, t+ = 0.5. For higher salt concentrations, the transference number decreases, reaching a value of approximately 0.2 for x ≥ 0.076. This value is in the range of the reported lithium transference numbers for standard PEO/LiClO4 complexes.35 The results confirm the predominant role played by the anion in the ionic conduction. The presence of excess lithium has little effect on the transference number, except for cation concentrations three times higher than the “mobile” anion concentration, where the overall conductivity is very low J



Article

composed of pendant methyl ether-polyethylene glycol chains and interconnected through polyether bridges. The dissociation state of the lithium salt was characterized by Raman spectroscopy, and the analysis showed that lithium perchlorate is completely dissociated up to a molar ratio x = 0.02, whereas, for higher molar ratios, the contact ion pair concentration increases linearly with the salt concentration. The materials strengths are the excellent thermomechanical stability and the high conductivity (up to 0.08 mS/cm at 30 °C). The electrochemical stability is also very promising, since the results suggest a possible application in combination with lithium metal anodes. Partial results from cyclization tests showed that Li/[HP/(LiClO4)0.05]/LFP cells can deliver 85− 90 mAh/g at 60 °C, 0.1 C, without showing any capacity fading during the first 20 cycles.

anyway. This is probably due to the scarce mobility of the excess lithium, which is “trapped” in silicate bonds. A complete analysis of the transport properties and of the conductivity mechanism will be the object of a future publication. Figure 15 shows the first 20 cycles of a Li/[HP/ (LiClO4)0.05]/LFP cell, being cycled at 60 °C and 0.1 C (17

■

ASSOCIATED CONTENT

S Supporting Information *

Additional information regarding GPC profile, FTIR, Raman, and NMR spectra of the precursor P, ATR-FTIR spectra, and conductivities of the seven. This material is available free of charge via the Internet at http://pubs.acs.org.

Figure 15. Charge capacities, discharge capacities, and Coulombic efficiencies of a Li/[HP/(LiClO4)0.05]/LFP cell, cycled between 3.0 and 4.0 V, at 60 °C and 0.1 C (17 mA/g, 0.04 mA/cm2).

■

mA/gLFP, 0.04 mA/cm2). Selected charge and discharge curves are shown in the Supporting Information (Figure S6). During the first ten cycles, the cell showed constant discharge capacity of 85 mAh/g, which is about 50% of the LFP theoretical capacity. Between the 10th and the 20th cycle, the discharge capacity increased linearly, reaching 90 mAh/g at the 20th cycle. The efficiency was also initially very low, about 70%, increasing progressively up to 95% after 20 cycles. The low discharge capacity (with respect to the theoretical value) is probably due to contact problems between cathode and electrolyte and more specifically to the difficult penetration of the electrolyte precursor in the cathode pores. This problem may be addressed either by optimization of the cell assembly, either by using the electrolyte as binder in the cathode formulation. The low efficiency values, observed during the first cycles, indicate that some irreversible process has been taking place in the system. However, the increasing efficiency and the absence of any discharge capacity fading suggest that these processes are to be ascribed to SEI formation and that they do not hinder the cell functionality (at least for the first cycles here examined).

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Author Contributions

All authors have given approval to the final version of the manuscript. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS We would like to thank Dr. Rüdiger Bertermann and Dr. Somchith Nique for taking the NMR spectra.

■

REFERENCES

(1) Di Noto, V.; Lavina, S.; Giffin, G. A.; Negro, E.; Scrosati, B. Electrochim. Acta 2011, 57 (SI), 4−13. (2) Bouchet, R.; Maria, S.; Meziane, R.; Aboulaich, A.; Lienafa, L.; Bonnet, J. P.; Phan, T. N. T.; Bertin, D.; Gigmes, D.; Devaux, D.; Denoyel, R.; Armand, M. Nat. Mater. 2013, 12, 452−457. (3) Wright, P. V. Br. Polym. J. 1975, 7, 319−327. (4) Goodenough, J. B.; Kim, Y. Chem. Mater. 2010, 22, 587−603. (5) Berthier, C.; Gorecki, W.; Minier, M.; Armand, M.; Chabagno, J. M.; Rigaud, P. Solid State Ionics 1983, 11, 91−95. (6) Ratner, M. A.; Shriver, D. F. Chem. Rev. 1988, 88, 109−124. (7) Doyle, M.; Fuller, T. F.; Newman, J. Electrochim. Acta 1994, 39, 2073−2081. (8) Gorecki, W.; Jeannin, M.; Belorizky, E.; Roux, C.; Armand, M. J. Phys.: Condens. Matter 1995, 7, 6823−6832. (9) Liang, S.; Choi, U. H.; Liu, W.; Runt, J. R.; Colby, H. Chem. Mater. 2012, 24, 2316−2323. (10) Schaefer, J. L.; Yanga, D. A.; Archer, L. A. Chem. Mater. 2013, 25, 834−839. (11) Bertasi, F.; Vezzù, K.; Negro, E.; Greenbaum, S.; Di Noto, V. Int. J. Hydrogen Energy 2014, 39, 2872−8883. (12) Stephan, A. M.; Nahm, K. S. Polymer 2006, 47, 5952−5964. (13) Weston, J. E.; Steele, B. C. H. Solid State Ionics 1982, 7, 75−79. (14) Kumar, B.; Rodrigues, S. J.; Scanlon, L. G. J. Electrochem. Soc. 2001, 148, A1191−A1195. (15) Croce, F.; Persi, L.; Scrosati, B.; Serraino-Fiory, F.; Plichta, E.; Hendikson, M. A. Electrochim. Acta 2001, 46, 2457−2461.

■

CONCLUSIONS We describe the synthesis, the structural properties, and the thermal and electrochemical stability of a series of 3D-hybrid inorganic−organic polymer electrolytes based on siloxane networks and cross-linked organic polar interphases. The materials were obtained by a two-step synthesis, the first step being a hydrolysis/co-condensation reaction between a PEGfunctionalized trimethoxysilane (K1) and a cross-linker (K2), a trimethoxysilane bearing an epoxy terminal group. Lithium hydroxide was added to the reaction mixture, in order to increase the structure disorder and to introduce some lithium in the co-condensate structure. The electrolytes were obtained by dissolving lithium perchlorate in the co-condensate and by cationic polymerization of the epoxide groups. The materials characterization, performed by vibrational and NMR spectroscopy, revealed that the electrolytes are composed by defect oligo-siloxane networks and by cross-linked organic interphases, K



Article

(16) Marcinek, M.; Bac, A.; Lipka, P.; Zaleska, A.; Zukowska, G.; Borkowska, R.; Wieczorek, W. J. Phys. Chem. B 2000, 104, 11088− 11093. (17) Kumar, B.; Rodrigues, S. J. J. Electrochem. Soc. 2001, 148, A1336−A1340. (18) Popall, M.; Du, X. M. Electrochim. Acta 1995, 40, 2305−2308. (19) Sanchez, C.; Belleville, P.; Popall, M.; Nicole, L. Chem. Soc. Rev. 2011, 40, 696−753. (20) Popall, M.; Andrei, M.; Kappel, J.; Kron, J.; Olma, K.; Olsowski, B. Electrochim. Acta 1998, 43, 1155−1161. (21) Skaarup, S.; West, K.; Zachau-Christiansen, B.; Popall, M.; Kappel, J.; Kron, J.; Eichinger, G.; Semrau, G. Electrochim. Acta 1998, 43, 1589−1592. (22) Popall, M.; Buestrich, R.; Semrau, G.; Eichinger, G.; Andrei, M.; Parker, W. O.; Skaarup, S.; West, K. Electrochim. Acta 2001, 46, 1499− 1508. (23) Bruce, P. G.; Vincent, C. A. J. Electroanal. Chem. 1987, 225, 1− 17. (24) Kim, G. T.; Appetecchi, G. B.; Carewska, M.; Joost, M.; Balducci, A.; Winter, M.; Passerini, S. J. Power Sources 2010, 195, 6130−6137. (25) Di Noto, V. J. Phys. Chem. B 2002, 106, 11139−11154. (26) Di Noto, V.; Vittadello, M.; Yoshida, K.; Lavina, S.; Negro, E.; Furukawa, T. Electrochim. Acta 2011, 57, 192−200. (27) Innocenzi, P.; Sassi, A.; Brusatin, G.; Guglielmi, M.; Favretto, D.; Bertani, R.; Venzo, A.; Babonneau, F. Chem. Mater. 2001, 13, 3635−3643. (28) Innocenzi, P.; Brusatin, G. Chem. Mater. 2000, 12, 3726−3732. (29) Di Noto, V.; Zago, V.; Biscazzo, S.; Vittadello, M. Electrochim. Acta 2003, 48, 541−554. (30) Di Noto, V.; Longo, D.; Münchow, V. J. Phys. Chem. B 1999, 103, 2636−2646. (31) Ducasse, L.; Dussauze, M.; Grondin, J.; Lassègues, J. C.; Naudin, C.; Servant, L. Phys. Chem. Chem. Phys. 2003, 5, 567−574. (32) Rosedale, J. H.; Bates, F. S. Macromolecules 1990, 23, 2329− 2338. (33) Thomas, K. E.; Sloop, S. E.; Kerr, J. B.; Newman, J. J. Power Sources 2000, 89, 132−138. (34) Chazalviel, J. N. Phys. Rev. A 1990, 42, 7355−7367. (35) Bouridah, A.; Dalard, F.; Deroo, D.; Armand, M. Solid State Ionics 1986, 18, 287−290.

L


Highly Conducting 3D-Hybrid Polymer Electrolytes for Lithium

Recommend Documents