Electrocatalysis Paradigm for Protection of Cathode Materials in High

Jul 6, 2016 - The cathode protection additives oxidize before the solvent and serve as sacrificial inhibitors of the catalytic centers. Without the ad...
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Electrocatalysis Paradigm for Protection of Cathode Materials in High-Voltage Lithium-Ion Batteries Ilya A. Shkrob* and Daniel P. Abraham Chemical Sciences and Engineering Division, Argonne National Laboratory, 9700 South Cass Avenue, Argonne, Illinois 60439, United States S Supporting Information *

ABSTRACT: A new mechanistic framework is suggested to account for the protective action of certain electrolyte additives on high-voltage positive electrode (cathode) materials. The mechanism involves inactivation of catalytically active centers on the electrode active materials through fragmentation reactions involving molecules at its surface. The cathode protection additives oxidize before the solvent and serve as sacrificial inhibitors of the catalytic centers. Without the additive, the surface oxidation of the solvent (like solvent oxidation in the bulk) yields H loss radicals and releases the proton that can combine with anions forming corrosive acids. This proton-release reaction is demonstrated experimentally for boronate additives. Specific radical reactions for the latter additives on the electrode surface are suggested. The same approach can be used to rationalize the protective action of other additives and account for various observations regarding their performance.

1. INTRODUCTION

coating the cathode with a thin layer of a less reactive material;50 it will not be considered here. While the development of high-voltage electrolytes and cathode protection additives is becoming a mature field, a concept that is comparable in its generality to Peled’s SEI is still lacking. In this study we aim to suggest such a holistic paradigm. Just like the concept of SEI, our conceptual framework does not seek to replace detailed understanding of the complex processes taking place on the cathode; rather, it provides a vista from which this complexity can be addressed. The article is organized as follows. In section 2 we briefly survey the current ideas regarding the mechanisms for cathode protection and solvent oxidation in LIBs and formulate the hypothesis (graphically sketched in Figure 1). In section 3 we provide details on experimental and computational methods used to generate the data. In section 4.1, we study the redox reactions of boronate additives (shown in scheme 1) using radiolysis and matrix isolation electron paramagnetic resonance (EPR). In section 4.2, we consider how our hypothesis can account for cathode protection by these additives. Section 5 presents our more general conclusions. To save space, some of the supporting schemes, tables, and figures have been placed in the Supporting Information. When referenced in the text, these materials have the designator “S”, as in Figure S1.

Nearly 40 years ago, Emanuel Peled introduced a new mechanistic paradigm in organic electrochemistrya solid− electrolyte interphase (SEI).1 This concept was suggested to explain the stability of carbonate electrolytes in the then nascent lithium ion batteries (LIBs), which contradicted the fact that these organic solvents readily reduced on Li metal and/or lithiated graphite (that serve as negative electrodes in LIBs). Peled realized that such a situation would be possible if a solid matrix consisting of electrolyte breakdown products on the electrode could pass Li+ ions through itself relatively easily while entirely blocking the access of solvent molecules to the reducing surface. This paradigm guided the subsequent development underpinning the commercial success of LIBs. Since its inception, much chemical detail was provided,2−9 but the original concept still informs our thinking about LIB operation. Presently there is another challenge before the scientific and engineering communities, viz. developing high energy density batteries for electric and hybrid vehicles, and once again LIBs can revolutionize the field. This application requires highcapacity (>220 mAh/g) cathode materials10 operating above 4.5 V vs Li/Li+.11−13 With such high-voltage cathode materials, protection of the solvent against anodic oxidation becomes a priority. Three general approaches have been pursued to address this concern. One approach is developing new solvents having greater stability on the cathode due to their higher oxidation potential (such as fluorinated carbonates,14−17 sulfones18−23 and sulfolanes,24 ionic liquids,25,26 and nitriles,27,28 etc.). The second approach is developing protective agents that serve as sacrificial additives (0.2−5 wt %) in regular carbonate solvents that decompose before the solvent and passivate the cathode surface.29−49 The third approach is © 2016 American Chemical Society

2. THE HYPOTHESIS Before proceeding further, we briefly examine the context. The high-energy cathodes currently come in two flavors: layered oxides (such as Mn- and Ni-substituted LiCoO2) that have lower operation voltages (∼4.4 V vs Li/Li+) but higher specific Received: June 7, 2016 Revised: June 30, 2016 Published: July 6, 2016 15119

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additive results in two (potentially related) desired outcomes: (i) preventing the corrosion and (ii) stopping the oxidative breakdown of the electrolyte. This passivation should not impede Li+ transfer, and the additive needs to be oxidized before the solvent. The most effective cathode protection additives typically are P- and B-containing compounds, including the boronates shown in Scheme 1. Two general ideas have been put forward regarding the protection mechanism. One idea is that the additives (or products of their oxidation, see below) react with HF and/or PF5 at the surface of the electrode, reducing the corrosion.45,50 The second idea is that oxidation of these additives yields a solid layer analogous to SEI, albeit much thinner, that somehow prevents these undesired outcomes.40,64,65 There is insufficient evidence for both of these rationales. In particular, there appears to be little correlation between the reactivity toward HF and the ability to protect the electrode,47 although in some cases there is definite connection.66 Furthermore, while it is natural to attribute cathode protection to the formation of a physical barrier (in analogy to SEI on graphite), it has proven surprisingly difficult to demonstrate that such a passivating barrier is indeed in place. A recent study67 indicated that trimethylboroxine additive (compound 3 in Scheme 1) yielded no solid deposits on the electrode that are detectable using Xray photoelectron spectroscopy. Another study43 indicated that BF3 was formed (as determined by head space analysis), suggesting that 3 reacted with HF. The authors suggested that BF3 promoted polymerization of the solvent with the formation of a protective layer, yet no evidence for this rationale was found. On the theoretical side, there are numerous studies of solvent oxidation,68−76 many of which share the same deficiency: the oxidation is considered as electron detachment from an isolated molecule (or a cluster of 1−2 such molecules and anions) in the gas phase. The ensuing reactions have little bearing on the actual redox chemistry in solution that is known, inter alia, from radiolysis of these solvents. The typical reaction is, in fact, deprotonation of the radical cation (SH+•) to another (or, sometimes, the same) solvent molecule (SH),74,77,78

Figure 1. Schematic representation of the electrocatalytic mechanism. When energized, the surface center on the lithiated transition metal (M) oxide deprotonates reaction 3) becomes an oxygen-center radical that abstracts H atom from the solvent (SH, reaction 4 that can combine with anions (A−) to yield corrosive acids. A sacrificial cathode protection additive (X−Y) inhibits the catalytic center by oxidative fragmentation as seen in reaction 5, as illustrated in the plot.

Scheme 1. Chemical Structures for Boronate-Based Cathode Protection Additives

capacities, and spinels (such as LiNi0.5Mn1.5O4) or phosphates (such as LiCoPO4) that operate at higher voltages (∼4.8 V vs Li/Li+), but have lower capacities.51 Delithiation of the layered oxides causes the formation of NiIV oxides that eliminate oxygen;52 this damage obviously cannot be prevented by electrolyte additives, although the effects of oxygen loss can in principle be mitigated. In both of these high-energy cathode materials, however, there is another type of damage that leads to irreversible loss of transition metal ions (e.g., Mn2+) to the solvent and their subsequent deposition on the negative electrode, where these ions have adverse effect on the cell performance (see ref 53 and references therein). Loss of the metal ions is explained by the formation of aggressive acids (such as HF), which degrade the material.54,55 The typical lithium salt in these electrolytes is LiPF6, as other salts cause dissolution and pitting of Al current collectors on which the cathode material is coated.6,56−60 The hexafluorophosphate anion is in equilibrium with PF5 and F−, reaction 155,61,62 LiPF6 ⇌ LiF + PF5

SH+• + SH → S• + SH 2+

(2)

resulting in the formation of a H loss radical and a protonated solvent molecule (in an electrolyte, the anion can be the preferred proton acceptor). For battery additives, similar reactions have been suggested by Borodin and co-workers76,79 from their extensive computational studies. While there are exceptions to reaction 2 (e.g., vinylene carbonate, see ref 74), these are uncommon, as in polar aliphatic solvents the intermolecular proton transfer is extremely efficient and rapid ( B−OH+−CH3 species). The resulting adduct SH2+ can recombine with the electrons e−• ejected from the parent molecule by ionizing radiation, yielding mobile H atoms that rapidly abstract H from another molecule: e−• + SH 2+ → H• + SH

(8)

H• + SH → H 2 + S•

(9)

Figure 2. First-derivative EPR spectrum observed from irradiated 4 at 50 K. The triplet of lines indicated by open circles originates from the H loss radical sketched in the inset. The outer lines (note the magnification factor of 20) originate from the dimer radical anion.

(Table S1), which are in good agreement with the experimental estimates. The inner resonance lines of this dimer radical anion overlap with the much stronger triplet (Figure 2) that originates from an •CH2O− radical that we identify with the H loss radical of 4 shown in Figure S2(b). As the temperature increases (Figure S4), the methylene group in this H loss radical begins to rotate more freely, and the hfc anisotropy becomes less pronounced, so the EPR spectrum better corresponds to the 1:2:1 triplet (expected in the fast rotation limit). The H loss radical and the dimer radical anion can be observed below 200 K; above this temperature the matrix softens and they begin to move and recombine. Thus, the redox reactions of 4 are fully analogous to 1. As shown below, this is, in fact, an exception rather than the rule: other molecules in Scheme 1 do not yield stable dimer radical anions. Even a compound that is very closely related to 1, trimethylboroxine (structure 3 in Scheme 1), yields a different kind of boron-centered radical. 4.1.3. Trimethylboroxine (3). In EPR spectra of irradiated 3, the dimer radical anion was not observed (Figure S5). Instead, there was a boron dangling bond center with the single boron11 nucleus having aiso = 248 G and B∥ = 15 ± 5 G. For the corresponding methyl loss >B• radical shown in Figure S1(c), our DFT calculation yields aiso = 256 G and B∥ = 20 G, which is in good agreement with the experimental observation. This suggests C−B bond scission. Indeed, in the EPR spectrum shown in Figure 3, trace i, there is a quartet separated by 22.3 G that corresponds to the methyl radical. At 150 K, this quartet disappears, while other radicals persist (Figure S6). As the methyl radical typically decays by H abstraction reaction 11

The excess electron can also recombine with the parent radical cation before it deprotonates via reaction 2, yielding a dissociative electronically excited state e−• + SH+• → SH* → fragmentation

•

CH3 + SH → S• + CH4

(10)

In matrices that readily trap electrons (like solid 1), reactions 8 and 10 are suppressed, whereas in matrices that do not readily trap electrons, these reactions become the main chemical pathways for excess electron decay, as discussed later. 4.1.2. Trimethoxyboroxine (4). In the EPR spectrum of irradiated 4 (Figure 2), there are resonance lines that can be attributed to a spin center (see Figure S3 and Table S1) with two magnetically equivalent 11B nuclei with aiso= 50.6 G and B∥ = 14 G. We identify this species with the dimer radical anion shown in Figure S1(b). Our DFT calculations yield a centrosymmetric structure 42−• with the B−B distance of 2.16 Å (vs 2.28 Å in 12−•) and aiso= 56.5 G and B∥ = 13 G

(11)

the residual radical is H loss radical of the parent compound (S•) shown in Figure S2(c). In Figure S7, we subtract the EPR spectra obtained at 50 K before (trace i) and after (trace ii) warming of the sample to 150 K (at which the methyl radical decays in reaction 11). Trace iii gives the difference trace that corresponds to the methyl radical, so the EPR spectrum shown in Figure 3, trace i originates through the superposition of EPR spectra from the methyl radical and the H loss radical generated via reaction 2. Subtracting the simulated EPR spectrum of the methyl radical (Figure 3, trace ii, we obtained the EPR spectrum shown in Figure 3, trace iii, which is very close to the EPR spectrum observed at 150 K (Figure 3, trace iv; however, 15122

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Figure 3. First-derivative EPR spectrum observed from irradiated 3 at 50 K (trace i). Trace ii is the numerical simulation of the EPR spectrum of the quartet of the resonance lines from the methyl radical (open circles). Trace iii is the difference trace. Trace iv is the normalized EPR spectrum from the same sample after warming to 150 K and subsequent cooling back to 50 K.

Figure 4. (a) First-derivative EPR spectrum observed from irradiated 5 at 50 K. The resonance lines indicated with the arrows are attributed to the H loss radical shown in the inset. The broad and poorly resolved singlet is attributed to 5−• also shown in the inset. (b) Temperature dependence of the EPR spectra for the sample shown above. As temperature increases, the H loss radicals decay, and the residual EPR spectrum is mainly from the more stable radical anion.

the latter still contains a small contribution from the methyl radicals. As suggested by EPR simulation in Figure S8 and comparison of the calculated and estimated hfc parameters in Table S1, this is indeed the >B−•CH2 radical (Figure S2(c)). It appears, therefore, that frozen 3 does not trap electrons: rather, the electrons decay via reactions 8 and 10, and the resulting excited state eliminates the methyl radical yielding the B-centered radical. 4.1.4. Phenyl Boronic Acid Ethylene Glycol Ester, PBE (5). PBE (structure 5 in Scheme 1) exhibits yet another mode of electron localization different from 1 and 4. The EPR spectrum was a superposition of the resonance lines attributed to the H loss radical (see Figure S2(d)) and another species, whose EPR spectrum corresponds to a broad singlet (Figure 4a). Just such an EPR spectrum is predicted for the monomer radical anion 5−• (Figure S9) stabilized through the electron delocalization in the aromatic ring (Figure S1(d)). The simulated EPR spectra for the H loss radical are given in Figures S9 and S10 and the parameters are given in Table S1. When the temperature increases (Figure 3b), this radical disappears from the EPR spectra, and only 5−• is observed. Thus, when the arene rings are involved, the formation of the monomer radical anions can be expected. 4.1.5. Tris(trimethylsiloxy)borate (2). The EPR spectra of irradiated 2 at 50 K are shown in Figures 5 and 6. While there are weak resonance lines from the B dangling bond center (Figure 5a), there are additional multiplets having the same hyperfine splitting as in the triplet shown separately in Figure 5b, suggesting that these features are silicon-29 satellites of the triplet progenitor. At 230 K (Figure 6), this triplet transforms into a septet with aiso = 4.8 G that is superimposed on a 1:2:1 triplet of narrow lines separated by ∼20 G. The latter are certainly from a •CH2− radical. The observed transformation of the anisotropic triplet into a septet is expected for a −•SiMe2 radical with the arrested rotation in the two methyl groups (Figure S11). In a conformationally locked radical shown in Figure S2(d) only two protons in these methyl groups are strongly coupled. At 230 K, there is rapid rotation averaging in

Figure 5. First-derivative EPR spectrum of irradiated 2 at 50 K. The central section (panel a) is comprised of overlapping EPR signals from the methyl loss (open circles) and H loss radicals (vertical arrows) shown in the insets. Panel b exhibits the wide-sweep EPR spectrum with the quartet of resonance lines from the >B• radical indicated with the open circles and the silicon-29 satellites of the methyl loss radical (with the calculated aiso(29Si) ≈ −177 G) indicated with the blue arrows. The doublet from the trapped H atoms is removed.

the radical, and the EPR spectrum corresponds to the septet (Figures 6 and S11). In addition to this septet there is a triplet from the − Si(Me)2•CH2 radical that yields the 1:2:1 triplet 15123

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Figure 7. Reaction scheme for inhibition of surface electrocatalysis by a boroxine compound reaction 14a. This sequence is a specific implementation of the general mechanism shown in Figure 1.

Figure 6. First-derivative EPR spectra observed from irradiated 2 as the temperature increases from 50 K (corresponding to the arrested rotation of the methyl and methylene groups in the radicals) to 230 K (corresponding to the free rotation). The resonance lines and features indicated with the open circles are attributed to the H loss radical, and the resonance lines indicated by the arrows are attributed to the methyl loss radical.

with the formation of an H loss radical (•R(−H)) in the arm, the boronate compounds can react in two other ways M−O• + BR3 → M−O−BR 2 + R•

above 200 K (Figure 6), when the methyl and methylene groups rapidly rotate (Figure S11). Therefore, in this system, in addition to the H loss radical, there is also a methyl loss radical. Could this radical be the product of one-electron oxidation? The asymmetric triplet shown in Figures 5b and 6 has been observed in radiolysis of hexamethyl disiloxane99 and other methylated siloxanes,100 and the resulting Si-centered radical can be identified using spin traps.101 Thus, the demethylation is the common reaction of such compounds, with the methane and H2 being the two principal products.99,100,102,103 Menhofer and Heusinger101 suggested that the approximate parity in the yields of the Me loss and H loss radicals originated through their common origin via reaction 2 that is followed by reaction 12 (instead of reaction 8) that is peculiar to these compounds −HO+ − SiMe3 + e−• → −O•SiMe2 + CH4

→M−O−B(R) R( −H) + RH

(14b)

Reaction 14b can occur if the leaving radical (•R) abstracts H atom from the bound fragment. Reactions 14a and/or 14b can be preferred to the competing H atom abstraction due to high energy of the resulting strong O−B bond. This preference can already be seen in the gas phase reactions of hydroxyl radical (Table 1), which is the simplest Table 1. Energetics of Gas Phase Oxidation of the Selected BR3 Compounds in Scheme 1 Involving a Reactive Oxygen Species (M−O• Radicals) Estimated Using the DFT Method reaction heat −ΔH, eV M in M−O•

BR3

leaving radical R•

reacn 13

reacn 14a

reacn 14b

H H H H H (HO)3Al− (HO)3Al− (HO)3Al− (HO)3Al− (HO)3Al−

1 2 3 4 6a 1 2 3 4 6

CH2OH CH2Si(Me)2OH CH3 CH2OH C6H5 CH2OH CH2Si(Me)2OH CH3 CH2OH C6H5

0.84 0.65 0.84 0.80 0.91 0.66 0.47 0.66 0.61 0.72

0.93 0.65 1.81 0.89 1.19 1.11 0.81 2.12 1.12 1.44

0.86 0.66 2.21 0.81 1.91 1.15 0.84 2.51 1.20 2.32

(12)

If this rationale is correct, oxidation of 2 is analogous to other compounds in Scheme 1 occurring via reaction 2. To summarize our EPR results, bulk oxidation of additives in Scheme 1 invariably yields H loss radicals via deprotonation reaction 2. Their one-electron reduction is inefficient unless there is the formation of a dimer radical anion (as for 1 and 4) or a monomer radical anion (as for 5). The latter requires the presence of the conjugated π-bond structure. All of the H loss radicals can readily abstract H from the carbonate solvent, suggesting that their oxidation in the bulk would give the same radicals as the direct oxidation of the solvent, resulting in a futile cycle. Thus, one-electron oxidation of the boronate additives in Scheme 1 cannot passivate the cathode any more than direct oxidation of the solvent. This paradoxical result hints that the oxidation occurring on the cathode must be unique to such energized surfaces (Figure 7). 4.2. The Mechanism for Electrode Passivation. Within our electrocatalysis framework (Section 2), the existence of the additional oxidation channel can be readily rationalized, as besides the H abstraction M−O• + BR3 → M−OH + •R( −H)BR 2

(14a)

•

a

Benzene ring addition of HO• radical is exothermic by 0.72 eV.

oxygen-centered radical of the postulated kind. According to our DFT calculations, the H abstraction by the hydroxyl from 4 is exothermic by 0.84 eV, whereas elimination of the hydroxymethyl radical via reaction 14a is exothermic by 0.93 eV. For 2 and 4, all three reactions have comparable exothermicity, for 3 and 6 reaction 14a is considerably more exergonic than reaction 13. Using (OH)3AlIII−OH− center as a d-ion model for the transition metal oxide, we observed similar bias toward the radical elimination reactions (Table 1). Once the M−OH centers are inactivated, the oxidation can only

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The Journal of Physical Chemistry C involve the bridging oxygen centers (M−O+•−M), and the rate of solvent decomposition decreases dramatically. Not only does this “electrocatalysis paradigm” allow rationalizing the protective action for the electrolyte additives, in some cases it can account for the solvent stability. For example, the current explanation for sulfone (RSO2R’) stability involves the effect of anion solvation on the oxidation potential of the molecule,72 although this conclusion seemingly contradicts the study of Borodin et al.74 indicating the facility of deprotonation. An alternative explanation is possible by using our hypothesis. It is known from the literature that oxidation of the sulfones by hydroxyl radicals yields unstable Hα loss radicals that undergo C−S bond scission reaction 15,104,105 e.g. HO• + RSO2 CH 2CH3 → H 2O + RSO2• + C2H4

very stable SEI on the graphite negative electrode and inactivate catalytic centers at the positive electrode. The quest for such additives continues both at Argonne and at research organizations worldwide.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpcc.6b05756. List of reactions and additional tables and figures (PDF)

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AUTHOR INFORMATION

Corresponding Author

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*(I.A.S.) Telephone: (630) 2529516. E-mail: [email protected].

One can envision the analogous reaction occurring at the surface and involving the M−O• center. The released radicals (RSO2•) cannot abstract H from other molecules; however, these radicals can recombine with the M−O• centers yielding alkylsulfonates

Notes

The authors declare no competing financial interest.

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yielding the same product. These reactions permanently inactivate the oxygen dangling bond center and account for the relative stability of the sulfones in the high-voltage cells; in this sense, this solvent is its own electrolyte additive.

ACKNOWLEDGMENTS This work was supported by the US-DOE Office of Science, Division of Chemical Sciences, Geosciences and Biosciences under Contract No. DE-AC02-06CH11357 to Argonne. D.P.A. is grateful for support from the US-DOE Office of Vehicle Technologies. The submitted manuscript has been created by UChicago Argonne, LLC, Operator of Argonne National Laboratory. The U.S. Government retains for itself, and others acting on its behalf, a paid-up nonexclusive, irrevocable worldwide license in said article to reproduce, prepare derivative works, distribute copies to the public, and perform publicly and display publicly, by or on behalf of the Government.

5. CONCLUSION In this study, we suggest a mechanistic framework to account how certain solutes can protect high-voltage cathodes. We suggest that this protection (Figure 1) involves inactivation of catalytic centers present on the energized transition metal oxide surface. That is, the electrolyte additives serve as sacrificial inhibitors of such catalytic centers. Without these additives, the oxidation at the electrode surface (like oxidation in the bulk) yields the corresponding H loss radicals and causes deprotonation of the solvent. The released protons combine with the anions present in the electrolyte to form corrosive acids damaging the electrode. The proton-release reaction 2 has been shown experimentally, for oxidation of the carbonate solvents90,92 and P- (e.g., refs 106 and 107) and B-containing additives (ref 98 and this study). Specific radical reactions for additives shown in Scheme 1 are suggested; such reactions for phosphite additives have been examined in ref 47. According to our view, sacrificial agents undergo oxidation fragmentation reaction 5 that inhibits the catalytic centers, so further oxidation can only occur at a higher potential. The deposition of solid residues and consumption of HF by the additives can be coincident with this underlying process. While our concept has been inspired mainly by photocatalysis, a rather similar set of ideas exists in aqueous electrocatalysis, where the corresponding reactions are known as Bockris’ electrochemical oxide path.108−110 Ending on a cautionary note, even the best performing additives at the positive electrode may, in fact, be inadequate for use in LIBs if they are unstable and reduce completely at the negative electrode. An ideal electrolyte additive would form a

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M−O• + RSO2• → M+−O3SR

(16)

Furthermore, direct reaction 17 (that is analogous to reaction 7 for phosphites) can occur for these M−O• centers only (as opposed to the hydroxyl radical), M−O• + RSO2 R′ → M+−O3SR′ + R•

(17)

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REFERENCES

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