Review: mechanisms and consequences of chemical cross-talk in advanced Li-ion batteries

Because H₂ and O₂ are known cross-talk agents in aqueous chemistries, gases are likely candidates for cross-talk phenomena in Li-ion batteries as well. During the formation of Li-ion batteries, gases are evolved as the electrodes react with the electrolyte and additives to form passivating layers on their surfaces. Many manufacturers remove these gases to prevent deformation of the cell, especially in pouch cells where the casing is flexible, and a uniform pressure is required during use. Dahn's group developed a device based on Archimedes' method to quantify gas production during battery operation and storage [17, 45]. Figure 2 shows an example of this measurement in Aiken's survey of high-voltage formation. Gas is generated during cell charging and consumed during storage [17].

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While it is difficult to determine the source of gas evolution in standard pouch cells, studies have been able to measure the high-potential evolution of H₂, CO₂, CO, and C₂H₄ via gas chromatography–mass spectroscopy (GC-MS) at each charged-state electrode by combining chemical analysis of both full pouch cells and charged individual electrodes, called 'pouch bags' [16, 46]. To assemble pouch bags, pouch cells were formed, degassed, charged, and cut open before their electrodes were separated and placed into individual sealed bags with either electrolyte or LiPF₆-free solvent. Non-degassed and degassed pouch cells were also studied as controls. To measure gas evolution during long storage periods, volume was measured via Archimedes' method and composition was measured by extracting the gas with a syringe and performing GC-MS [16], or through a gas extraction device which was inserted directly into a GC [46] that could detect H₂, CO₂, CO, and light hydrocarbons. To measure consumption of gaseous products, a known volume of gas was inserted into pouch bags containing a single charged electrode, and volume change during storage was measured using Archimedes' method [46]. In non-degassed full cells, the gases were consumed over time [46].

On-line electrochemical mass spectrometry (OEMS) or related differential electrochemical mass spectrometry (DEMS) can also be used to determine the impact and extent of cross-talk. For example, experiments show more electrolyte oxidation products are produced at the cathode when the anode is LiFePO₄ (LFP) instead of Li metal or graphite, because the low potential scavenges oxidation products [47]. Metzger et al proposed a two-compartment OEMS cell to rigorously separate the effects of cross-talk and to measure real-time gas composition at each electrode in operando [48]. In the design (figure 3), a Li-ion conductive glass ceramic (LiCGC) separates cells into two compartments, so gases and liquid-phase products evolved at one electrode may not interact with the other electrode [48, 49]. OEMS/DEMS also allows isotopic labeling experiments to determine the source of carbon, oxygen, and other elements in various decomposition products.

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Several transition metal oxide chemistries, including LMO, LNMO and various NMC forms (Li_xNi_aMn_bCo_1-a-b O₂, Li- and Mn-rich xLi₂MnO₃ · (1-x)LiMO₂ [M = Mn, Ni, or Co]), are known to suffer from metal dissolution. Operation at elevated temperature often exacerbates this phenomenon, while increasing the upper cut-off voltage during cycling has been shown to accelerate metal dissolution [58, 69–76]. Additionally, electrolyte identity can strongly influence metal dissolution, with fluorinated salts, particularly LiPF₆, promoting dissolution due to HF impurities [77]. While the trends of dissolution with voltage, temperature, and salt are clear, the mechanism by which metals dissolve out of these active materials remains an area of active research.

The earliest works aiming to understand metal dissolution from transition metal oxide active materials were investigating LMO-type electrodes. Thackeray and colleagues attributed Mn dissolution from LMO [78] to the Hunter disproportionation reaction [79], whereby surface and bulk Mn³⁺ react to form bulk Mn⁴⁺ and a surface Mn²⁺ which dissolves into solution. Aurbach proposed another mechanism of high-potential dissolution caused by HF corrosion [80], where electrolyte oxidation induced high Lewis acid concentrations that catalyzed dissolution and passivated the surface with LiF and MnF₂. Amatucci and colleagues found that the concentrations of dissolved Mn and HF increased proportionally when LMO electrodes were left submerged in electrolyte for several weeks [81]. Blyr similarly showed that intentional electrolyte contamination by water increased the extent of Mn dissolution from LMO [77], and both Blyr and Amatucci clearly demonstrated that storage at elevated temperatures exacerbated this dissolution [77, 81]. Blyr also showed that Mn loss was greatest when LMO was stored in LiPF₆-containing electrolyte compared to other fluorinated salts such as LiBF₄ and LiAsF₆, indicating that the instability of the PF₆^- anion plays an important role in Mn dissolution.

Terada used total reflection x-ray fluorescence to track electrolyte Mn concentrations [82], and found the oxidation state of dissolved Mn to be Mn²⁺. Recently, Banerjee used electronic paramagnetic resonance (EPR) paired with inductively coupled plasma spectroscopy and x-ray absorption near-edge spectroscopy (XANES) to determine the oxidation state of dissolved Mn in an LMO–graphite battery [83]. They found that dissolved Mn existed mostly in the Mn³⁺ state, with an average oxidation state of 2.6 ± 0.2, disputing early assumptions of dissolved Mn²⁺, and providing evidence for an alternative mechanism of Mn dissolution. Phase transformation at high voltages involving the formation of a soluble Mn₃O₄ phase could account for Banerjee's findings, since this mechanism involves direct dissolution of Mn(III) and Mn(II). Tang used scanning transmission electron microscopy to confirm the presence of Mn₃O₄ on highly charged LMO, and found the phase disappeared or dissolved during discharge [84]. An important distinction between the phase transformation and disproportionation mechanisms is not only the valence state of dissolved Mn, but also the concomitant oxygen release under phase transformation. While the disproportionation mechanism was extensively considered in early works, it does not fully encompass all modes of metal dissolution such as at high voltages or for Ni and Co dissolution from LNMO and NMC chemistries. Thus, disproportionation of Mn³⁺ may contribute to metal dissolution from high-energy cathodes, but it is neither the sole mechanism nor the most deleterious.

Similar mechanisms have been considered to better understand metal dissolution in other high-energy cathode materials. For instance, the disproportionation mechanism was later applied by Gallus to explain metal dissolution from NMC111 at low states of charge, while also considering HF corrosion as the cause of degradation at high states of charge [73]. Other proposed mechanisms involve dissolved organic species facilitating metal dissolution by forming soluble chelation complexes with those metals, such as fluorescent acetylacetonate complexes formed at LNMO [63, 85]. Lastly, researchers have shown that some cathode degradation pathways in layered NMC electrodes, such as oxygen release [84, 86–89], do not directly involve metal dissolution but rather create conditions favorable to dissolution on subsequent cycles. For instance, oxygen loss is thought to be accompanied by surface reconstruction and vacancy evolution, encouraging transition metal motion from bulk sites to surface sites, causing surface instabilities, and ultimately leading to the dissolution of the unstable, reduced transition metal surface layers [67, 88, 90, 91]. Efforts to impede metal dissolution primarily focus on coatings, which may increase impedance, or additives, which must not introduce new detrimental effects at the anode.

While the prevailing consensus is that dissolved Mn is present as Mn³⁺ or Mn²⁺, there are few reports on the oxidation states of dissolved Ni and Co and the assumption of Ni²⁺ and Co²⁺ has been largely unchallenged [64, 92]. Additionally, mitigation efforts to scavenge or trap metals before reaching the SEI are hampered by uncertainty over the transport mechanism of dissolved metals. The simplest explanation for metal transport is a migration mechanism in which charged Mn, Ni, and Co ions travel along electric field paths to the anode (figure 7(a)). However, this interpretation breaks down when considering the relative concentrations of dissolved metals. Terada measured Mn concentrations of ∼10 ppm for LMO cycled at room temperature and ∼400 ppm at 50 °C [82], while Amine more recently measured <20 ppm each for Mn, Ni, and Co dissolved from a Li_1.1(Ni_1/3Mn_1/3Co_1/3)_0.9O₂ electrode left soaking at 55 °C for one month [93]. Lee summarized the work of many groups and found that 50–200 ppm was a reasonable concentration range of dissolved metals for model and theoretical studies [94]. Clearly, these metals are extremely diluted when compared to conductive salts at concentrations of 1 M or greater. Thus, dissolved transition metal ions are effectively in the supporting electrolyte, and their transport is governed by diffusion alone.

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Shkrob originally suggested that electrolyte degradation products, such as oxalate, carbonate, or ethylene dicarbonate, form stable complexes with Mn²⁺, thus facilitating metal dissolution from the cathode [95]. They demonstrated the preferential deposition of Mn on mineralized SEI components and found through electron spin echo envelope modulation that Mn²⁺ was always coordinated to an organic ligand. These findings further discredit a transport mechanism based on migration of charged metal species. Shkrob then considered an alternate transport mechanism, shown in figure 7(b), in which metal ions are coordinated by electrolyte decomposition products such as Li ethylene dicarbonate (LiEDC). They reasoned that these decomposition products must be stronger chelating agents than outer SEI organic species, since only the strongly chelating carbonates present in the inner SEI layers can supplant the proposed chelating agents that facilitate metal dissolution at the cathode. Jarry and colleagues later used findings from an x-ray fluorescence study to propose that electrolyte oxidation at LNMO generates β-diketonate ligands which chelate surface metals and together dissolve as a neutral complex, and travel to the anode in this form before deposition (figure 8) [63].

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Tornheim and colleagues recently provided more evidence for this complexing mechanism by performing carefully designed pre-formation experiments [96]. Their work suggests that reduction products generated at the anode not only play a role in metal dissolution at the cathode, but also that metal penetration into the SEI is dependent on the chemical nature of the electrolyte used to pre-form the SEI. In fact, anode-dependent impedance rise at the cathode may be understood to involve this exact cross-talk mechanism. Rodrigues demonstrated that metal deposition rates on graphite were almost double those for the 'SEI-free' LTO [66], which may be indirect evidence that metal dissolution is inhibited by pairing cathodes with higher-potential anodes that do not generate the same chelation agent byproducts as graphite.

In a more recent communication, Tornheim showed that the coordinating ion of dissolved Mn strongly impacts cell cycling behavior: PF₆^- -coordinated Mn showed persistent side reactions and constantly decreasing CE, while TFSI^- -coordinated Mn had a large initial capacity loss, but stable CE during cycling (figure 9) [98]. Wang's recent communication complements Tornheim's studies, using theoretical interaction energy calculations and storage tests to determine that dissolved Mn²⁺ is preferentially solvated by TFSI^- over PF₆^-, yet the coexistence of PF₆^- and Mn²⁺ in electrolyte results in the greatest destabilization and degradation of the electrolyte [97]. This may explain why Tornheim observed that Mn(TFSI)₂-doped PF₆^–-containing electrolyte does not show the same performance as Mn(PF₆)₂-doped electrolyte: Mn²⁺ mostly remains solvated by TFSI^- anions, whereas the PF₆^- originating from the LiPF₆ salt induces similar side reactions to if no Mn salt were present (i.e. CE match between Gen2 and Mn(TFSI)₂ in figure 9(b)).

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Perhaps more importantly, Wang demonstrated that Mn²⁺ is preferentially solvated over Li⁺ by EC, TFSI^-, PF₆^-, etc, and concluded that even before reaching the SEI, dissolved Mn parasitically destabilizes solvation sheath members for chemical decomposition and activates electrolyte components for electrochemical reduction at the anode. This indicates that mitigation strategies based on SEI engineering may not completely prevent metal-induced parasitic reactions and overlook a mode of dissolved metal-induced capacity fade. These findings further corroborate Shkrob's proposed mechanism of metal transport between electrodes and underscore the importance of understanding the chemical nature of dissolved metals, since the identity of metal-coordinating species corresponds to different degradation pathways and battery lifetime outcomes.

Tornheim's and Wang's results taken together demonstrate that the solvation environment of metals within the electrolyte governs their deposition at the anode [96, 97, 98]. It is clear that neither electrochemical reduction to a metallic state nor ion exchange with –CO₂^-Li⁺ members of the organic SEI layer are occurring, due to the predominant location of Mn²⁺ on inorganic Li₂CO₃ crystallites [95]. Rather, metals are incorporated in the SEI layers by a mechanism involving their chelating agents, and their destination within the SEI is governed by the chemical character of the SEI. For instance, Mn is commonly observed near the surface of an SEI formed in highly fluorinated electrolyte, whereas higher bulk Mn concentrations are observed for SEI formed in conventional electrolyte [96]. Once trapped in the SEI, these metals destabilize the protective layer leading to unabated capacity fade.

The main source of capacity fade in high-energy Li-ion batteries originates from metal contamination of the SEI, yet there is no generally agreed-upon mechanism of SEI failure. Delacourt suggested that the SEI's insulating properties are circumvented by metallic Mn deposits forming an electronically conductive pathway (figure 10(b)) [60], which Gowda later supported based on their finding of Mn(0) in the SEI of charged graphite with XANES [59]. Delacourt suggested another possibility for Mn-induced capacity fade involving the formation of a highly porous SEI that allowed facile electrolyte transport to the anode for reduction (figure 10(c)) [60]. Ochida similarly reported that Mn particles decompose surface films (i.e. SEI), and Joshi specifically ascribed metal-induced capacity fade to the decomposition of LiEDC, revealing fresh graphite surfaces for electrolyte reduction [8, 68].

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Joshi also noted that Mn promotes the formation of inorganic LiF and Li₂CO₃ species within the SEI, as products of LiEDC decomposition [8]. Shkrob used extended x-ray absorption fine structure and EPR to show that the SEI contains a reduced form of Mn below the +2 valence state, but not metallic α-Mn(0) [95]. Shkrob interpreted this finding as evidence for Mn acting as a catalytic center for SEI decomposition. Vissers found that inorganic SEI species are able to sequester Mn ions [99], and supported Shkrob's proposed mechanism by asserting that Mn is reduced to the +1 oxidation state during graphite lithiation and subsequently oxidizes to +2 by catalytically reducing adjacent electrolyte molecules. Wang found evidence supporting an electrocatalytic reduction cycle involving Mn²⁺, and showed that Mn²⁺ is more effective than Mn⁰ at inducing electron transfer from the anode to facilitate EC reduction [97]. Based on calculated electron affinities of electrolyte reduction reactions, Wang asserted that single electron reduction of a solvated Mn²⁺ is far more likely than two-electron reduction [97].

Gilbert combined insights from Leung [100] and agreed that metal-facilitated capacity fade occurs via an electrocatalytic process of metal reduction and decomposition of organic SEI components (figure 11) [75]. More recent work from the Gasteiger group has provided support for this electrocatalytic metal redox mechanism of SEI decomposition, while also linking metal contamination to gassing events in Li-ion batteries [61, 101]. Our own recent efforts have found that predicted voltammetric curves of the electrocatalytic mechanism are quantitatively consistent with experimental measurements. However, the trends predicted by the thermodynamic redox potentials of Mn, Ni, and Co were obscured by kinetic differences between the metals. These findings further emphasize the importance of local coordination environment for metals within the SEI [56, 102, 103]. Understanding how ligands and anions affect the location and reactivity of metals within the SEI will be necessary to design mitigation approaches.

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Perhaps the most effective single approach to mitigating the effects of metal dissolution is engineering a more robust SEI. However, preventing metal incorporation into the SEI with SEI additives or active-trap separators does not prevent dissolved metals from destabilizing electrolyte components and consuming cyclable Li. The literature thus suggests that only a strategy based on eliminating dissolution can achieve high CE and cell lifetimes while maintaining the improved energy and power densities that motivate the use of these cathodes. As such, SEI engineering can help limit metal dissolution by preventing generation of soluble higher-molecular-weight electrolyte degradation products. These products can be particularly detrimental if they promote transition metal dissolution by acting as chelating agents or generating protons during cross-talk oxidation. Therefore, careful design of the SEI to avoid soluble products can impede cathode degradation and diminish the extent of metal dissolution, but it cannot fully eliminate this phenomenon. To move forward, researchers must focus on low-impedance coatings that stabilize the cathode surface or CEI-formation additives that neutralize the various mechanisms of metal dissolution. Additives that promote Li₂O, LiF, and Li₂CO₃ may be of particular interest, as these species are understood to provide important passivation at the anode. However, care must be taken to ensure that additive decomposition does not initiate new, harmful cross-talk mechanisms. Ultimately, continued investigation into how the CEI regulates metal dissolution will better inform advanced cathode designs for next-generation Li-ion batteries.

At thermodynamic equilibrium, the carbon and oxygen atoms of organic electrolytes will be converted into CO₂. However, a 5 V battery operates very far from equilibrium. It is thus not surprising that CO₂ is both generated and consumed by a number of heterogeneous and homogeneous processes throughout a battery. The most obvious source of CO₂ in a Li-ion battery is electrochemical oxidation of the organic electrolyte. Early voltammetry studies of a variety of carbonate electrolytes reported anodic stability limits between 3.3 and 4.3 V on planar gold, aluminum, and platinum electrodes [104, 105]. It was assumed that the current was due to solvent and salt oxidation with products of CO₂, CO, and aldehydes. However, this interpretation is inconsistent with the performance of commercial batteries, as noted by Jung et al [106]. At a typical commercial electrode loading of 10 mg cm⁻² oxide and 0.2 mg cm⁻² conductive carbon with specific surface areas 0.25 and 65 m² g⁻¹, respectively, Moshkovich's measurements at 4.5 V predict a parasitic current due to solvent oxidation that is at least 100 times the C-rate. Because of this complexity, direct detection of gaseous products using OEMS and DEMS is generally accepted as necessary to deconvolute solvent decomposition from other oxidation processes. Even then, there are a variety of reports in the literature regarding the potential at which electrolyte degradation occurs and its surface sensitivity, as discussed below. CO₂ and, to a lesser degree, CO are evolved at potentials as low as 4.2 V vs. Li, depending on the exact system. For many systems, almost all the gas in the first charging cycle can be attributed to the oxidation of residual carbonate impurities on particle surfaces [107]. Direct electrolyte oxidation is then not observed until almost 4.8 V. Many electrolyte additives are also preferentially oxidized at relatively low potentials [108], and CO₂ generation from this process has been directly observed.

The surface sensitivity of electrochemical solvent oxidation reactions is important for interface design but remains a subject of open debate in the literature. Certain cathode materials are clearly more active for CO₂ and O₂ evolution than others. Chemical intuition additionally suggests that electrocatalysis plays a role because solvent oxidation requires multiple electron and/or proton transfers. Multi-electron reactions are most commonly inner-sphere reactions that proceed through several adsorbed intermediates. Because the free energy of adsorption is surface-specific, this would indicate a surface-sensitive or electrocatalytic reaction. Density functional theory calculations predict that carbonate adsorption and dissociation is preferred on oxides with lower p-band centers [109], and other evidence suggests that electrolyte preferentially oxidizes at oxide vacancies or defects [63, 110]. Computational evidence against electrocatalytic solvent oxidation proposes that organic carbonate oxidation is initiated by removal of one electron to form a cation radical [111, 112]. This outer-sphere reaction would likely be much less sensitive to electrode surface structure. Experimentally, Jung et al showed that, after normalizing for surface area, C65 carbon and LNMO had the same profiles for current and CO₂ evolution with potential (figure 12). While layered NMC622 (Li[Ni_0.6Mn_0.2Co_0.2]O₂), NMC111, and NMC811 (Li[Ni_0.8Mn_0.1Co_0.1]O₂) had lower onset potentials for current and CO₂ and CO evolution, the gases were shown to originate from homogeneous reaction of solvents and electrochemical generation [106]. Metzger et al determined via isotopic labeling that a majority of CO₂/CO evolution occurs from electrolyte oxidation, while the remainder comes from electrode oxidation [113]. Furthermore, the electrolyte is more sensitive to oxidation at elevated temperatures than the electrode.

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Homogeneous electrolyte decomposition initiated by electrochemically generated species explains much of the variation between electrode materials. Oxygen can be released from metal oxides at very low levels of lithiation or high state of charge. While the state of charge corresponds with potential, oxygen release is often driven by structural instability, rather than electronic driving force. Wandt et al and Mahne et al showed that singlet O₂ was released from both active materials and Li₂CO₃ [114, 115]. Castel et al suggested that trace water could result in a variety of reactive oxygen species [116]. The subsequent reactions between reactive oxygen and organic electrolytes generate CO₂ and other products, which can then diffuse to the anode. The possibility that reactive oxygen remains stable long enough to react heterogeneously at the anode is discussed in section 3.3.

CO₂ can also be generated thermally at elevated temperatures. Sloop et al showed using thermogravimetric and chemical analysis that LiPF₆ decomposes into LiF and PF₅ at temperatures as low as 50 °C [14]. As a strong Lewis acid, PF₅ reacts with EC to form oligomers and CO₂. Xiong et al observed more CO₂ evolution in a pouch bag with a delithiated cathode than in a full pouch cell and no CO₂ evolution in a bag with a charged anode [16]. There was also less CO₂ evolution with no LiPF₆ in the solvent [16]. A number of other studies have reviewed the thermal decomposition reactions of salts and solvents [108, 117, 118].

The presence of CO₂ in Li-ion cells seems, from most reports, to be strongly beneficial to cell lifetime. CO₂ was an early electrolyte additive; Plichta et al attributed improved cycle life of LiCoO₂ –Li to reaction between CO₂ and trace LiOH and Li₂O cells to form Li₂CO₃ [119]. Aurbach found similar results, and speculated that trace CO₂, introduced from decomposition during solvent purification, could be a source of irreproducibility between research groups [120]. More recently, Krause found greatly improved lifetime on a Si anode when a small amount of dry ice was incorporated into the cell [121].

Because LCL at the anode dominates most capacity fade, the mechanism of successful electrolyte additives is most likely improving the passivity of the SEI. Lucht et al showed that in the absence of CO₂, SEI chemistry contained very little Li₂CO₃ [122, 123]. Electrochemical reduction of CO₂ in aprotic media was studied extensively by Saveant's research group and concluded that on inert electrodes such as mercury and lead, reaction proceeded through the CO₂ radical anion [124–126]. Conductive carbon such as graphite is also generally considered catalytically inert and would likely follow a similar reaction scheme. Schwenke et al completed a rigorous OEMS study of CO₂ at low potentials [127]. Flowing isotopically labeled [13]CO₂ into the OEMS cell showed that CO₂ is not reduced to CO. Instead, it is consumed primarily by three processes, depicted in figure 13. At 1.5 V, electrochemical reduction of trace water increases the pH and converts CO₂ to CO₃^-, which then precipitates as Li₂CO₃. At 0.7 V, CO₂ is reduced directly to CO₂^-·. Following the findings of Saveant [124–126], the radical can combine with trace protons to formate or can dimerize to oxalate. Finally, at 0.01 V, CO₂ can be reduced to solid amorphous carbon and carbonate. Carbonate and oxalate have very low solubility in the electrolyte, and thus form more passivating layers than organic SEI products, thereby explaining the beneficial effects of CO₂ on cycle life.

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While CO₂ has generally been shown to benefit cell lifetime, CO₂^-· may also be detrimental, as it is highly unstable and can react homogeneously with solvent as well [100]. Tang et al and Ma et al used an interesting cell design to flow gases while cycling a cell and found that the capacity fade in several cell chemistries was twice as rapid when cycled in 1% CO₂ in Ar vs. pure Ar [130, 131]. The net benefit of CO₂ in a battery will certainly depend on the relative rates of dimerization, precipitation, and further reaction with solvent, all of which are likely sensitive to local transport and electrode geometry as well as surface chemistry.

Along with CO₂, CO is also typically generated in electrolyte degradation processes, but at lower concentrations. There is very little research on the direct effects of CO. Aqueous electrochemical research on CO₂ reduction shows that CO₂ and CO reduction often yield similar end products. If this trend applies in nonaqueous batteries as well, CO reduction is likely also beneficial to anode passivation. However, unlike aqueous electrochemistry, CO₂ reduction in battery solvents does not yield CO as a product [127], showing that more work is necessary to understand this family of reactions.

The SEI formation process is known to result in gases including H₂, CO₂, CO, CH₄, C₂H₄, and C₂H₆. Gas generation is known to increase with higher water content, which can happen due to contamination during cell assembly or when electrodes are poorly dried. Bernhard et al used OEMS with one- and two-compartment cells to investigate one-electron reduction of water at the anode forming H₂ and OH^- [49]. They proposed that the OH^- triggers a ring-opening reaction of cyclic carbonates, causing solvent decomposition into CO₂. Metzger and colleagues used the same OEMS setup to observe that without H₂O contamination, the source of H₂ evolution is due to protic electrolyte oxidation species (R-H⁺) that diffuse from the cathode and subsequently are reduced at the anode [48].

Unlike O₂ and CO₂, which are generated and consumed at different electrodes, reduction gases seem to be both generated and consumed primarily at the anode. Ellis et al injected pure H₂, CO₂, CO, C₂H₄, CH₄, or C₂H₆ into pouch bags with only one electrode in order to determine the reactivity of gases at each electrode [46]. Large volumes of CO₂, CO, H₂, and C₂H₄ were consumed at the anode, presumably by the mechanisms discussed in section 3.2. At the cathode, H₂ was partially consumed, and although no volume change was observed with C₂H₄, a large charge transfer resistance increase was attributed to C₂H₄ oxidation. No saturated hydrocarbons were consumed in either case. In full cells, where the anode was able to consume evolved gases, no adverse effects on cell performance were observed [46]. There has been very little investigation on direct effects of gases other than CO₂, but one study found that CH₄ on its own can increase capacity fade, but is counteracted by CO₂, so that when both are present in the cell no capacity fade is observed [130]. A specific mechanism is uncertain; however, alkanes formed at the anode during SEI formation can be oxidized at the cathode, generating protic species (as discussed in 4.1.1) while CO₂ has been shown to act as a proton scavenger (as discussed in 3.2.2) reacting with protic species to generate formate and passivate additional reactivity [127, 142].

For gases produced at the anode, cross-talk to the cathode seems to be generally detrimental. Rodrigues found that if gases were not consumed by the graphite anode, their diffusion to the cathode caused an impedance increase [66]. However, it is not clear that gases, rather than higher-MW products (section 4.2), were responsible for the impedance rise. Venting the cell after formation cycles can significantly reduce capacity fade, pointing to the detrimental long-term role of gas chemistry. H₂ gas formed at the anode has also been reported to accelerate capacity fade and increase impedance at the cathode in graphite//LCO cells [130]. Metzger argued that electrochemical H₂ oxidation at the cathode forms protons that react with PF₆^- anions to form HF. HF can react with Li salts to precipitate as insulating LiF. HF also accelerates metal dissolution and associated problems (section 3.1.2). In the same work, electrode cross-talk was shown to accelerate H₂ production [48]. NMC/graphite full cells showed higher H₂ product relative to graphite half-cells due to diffusion of protonated electrolyte oxidation products from the cathode to the anode [48]. H₂ has also been shown to increase cracking in the cathode [66, 130]. H₂ is known to embrittle steel and other materials, and Rodrigues et al proposed a similar mechanism via crack propagation induced by local changes to the plasticity and critical stress. Ma proposed a proton-induced decrease in the cohesive strength of O₂^- sites within the oxide leading to crack formation and mechanical failure.