The Electrochemical Mechanisms of Solid–Electrolyte Interphase

Aug 8, 2019 - In this work, we adapt methods from electrocatalysis to allow for the potential-dependent, atomistic simulation of SEI formation on lith...
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The Electrochemical Mechanisms of Solid−Electrolyte Interphase Formation in Lithium-Based Batteries Martha A. Gialampouki,‡ Javad Hashemi,‡ and Andrew A. Peterson*

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School of Engineering, Brown University, Providence, Rhode Island 02912, United States ABSTRACT: In lithium-based batteries, the solid−electrolyte interphase (SEI) is a layer of material that forms between the negative electrode and the liquid electrolyte; it is produced spontaneously by the breakdown of electrolyte compounds at the highly reducing potentials inherent to these systems. The SEI is perhaps the most important factor controlling the efficiency, safety, and lifetime of lithium batteries, and many empirical approaches have been developed to control the SEI’s properties. In this work, we adapt methods from electrocatalysis to allow for the potentialdependent, atomistic simulation of SEI formation on lithium surfaces via electronic structure calculations. We use a computational lithium electrode (CLE) technique, in which the potential scale is thermodynamically linked to the lithium reference electrode, to study the decomposition of ethylene carbonate, one of the most prevalent battery electrolytes, into SEI components on lithium-based surfaces. On Li metal surfaces, we find that the most favorable process is forming the inorganic carbonate phases (accompanied by the liberation of ethylene gas), while forming the organic SEI component lithium ethylene dicarbonate (LiEDC) is unfavorable. In contrast, we find LiEDC to be favorable on the inorganic lithium surfaces (Li2CO3 and Li2O). This gives a mechanistic interpretation of a common physical picture of SEI formation: inorganic species (e.g., Li2CO3) are formed more heavily near the electrode, while organic species (e.g., LiEDC) are formed more heavily near the electrolyte. Both electrochemical and non-electrochemical pathways are explored and found to have similar energetics at this level of theory. This study rationalizes experimental findings and sets the stage for mechanism-based control of SEI formation.

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equilibrium potential of Li(s) ↔ Li+ + e−, which is in equilibrium at −3.04 V versus the standard hydrogen electrode at room temperature. This creates a challenge in finding electrolytic solvents that are stable under such strong reducing conditions, and in practice, the electrolytes employed in modern batteries, such as ethylene carbonate (EC), are reduced via cathodic reactions at this electrode. Fortunately, the breakdown products of this reduction form a passivating layer on top of the electrode, further slowing solvent breakdown; this passivating layer is referred to as the solid− electrolyte interphase (SEI). 7,21−25 If this material is conductive to Li+ but electrically insulating, then normal battery operation can still take place. A stable SEI can also enhance the batteries’ safety by protecting the anodes from Li dendrite growth on Li-metal or continuous graphite exfoliation on Li-ion batteries.8,26−28 As the SEI increases in thickness, it acts as a stronger resistor, reducing the operating efficiency of the battery by generating heat upon charging and discharging7,21,29−34 and consequentially decreasing the system’s safety and lifetime.7,22−24 Further, as the SEI grows, much of the reactive Li of the battery is trapped in this layer, also reducing

INTRODUCTION Energy storage is among the primary industrial needs and scientific challenges of the modern age, and lithium-based batteries are prominent candidates to address this challenge. Lithium-ion batteries (LIBs) are currently the most-used type of batteries in portable devices, from small electronics to electric vehicles and aerospace,1−7 due to their energy densities and theoretical areal capacity, which are crucial for their longrange performance.7 On the other hand, the increasing demand for higher energy-density batteries has encouraged the scientific community to re-examine the feasibility of safe rechargeable lithium metal batteries.8,9 Several groups have developed fabrication techniques leading to dendrite-free lithium deposition,10−13 while other efforts have been devoted to more alternative approaches, such as solid-state or polymer electrolytes,14−16 mechanical modification of Li surfaces,17 and controlled lithium solvation.18,19 All conventional lithiumbased batteries consist of three essential components: the negative electrode (typically graphite or metallic Li), the positive electrode (typically a metal oxide in current batteries, e.g., LiCoO2, LiMn2O4, etc.), and an electrolyte to provide electrical insulation while allowing Li+ mobility (typically a mobile lithium salt dissolved in an organic solvent).20 The negative terminal of such batteries is extremely reducing: the electrode operates at potentials near the © XXXX American Chemical Society

Received: April 25, 2019 Revised: July 3, 2019

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DOI: 10.1021/acs.jpcc.9b03886 J. Phys. Chem. C XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry C

of the SEI synthesis, which ultimately can lead to improvements in existing batteries and new electrode materials.

the capacity of the battery. Thus, the SEI is a primary contributor to long-term capacity fade, efficiency losses, and side effects such as resistive heat generation, but its presence is essential to the stability and safety of such batteries. Thus, an ideal SEI is one that can be kept thin and stable, and much work has focused on this goal.7,21,29−34 The SEI is a heterogeneous, largely amorphous layer of organic and inorganic components.7,21 The properties and composition of SEIs have been widely studied both experimentally and theoretically.7,30,31,33,35−41 Experimental studies have detected several compounds in the SEI such as lithium ethylene dicarbonate (LiEDC), Li2O, Li2CO3, LiOH, CH3OLi, CH3OCO2Li, and LiF. Among them, LiEDC has been heavily studied as one of the most common organic SEI products.27,42−44 Different SEI formation mechanisms have been proposed, where the most common are the interfacial reactions between the positively charged lithium ions, (Li+), the organic solvents of the electrolyte (e.g., EC), and the electrons from the negative electrode.21,22,31,36,40,42,45−47 These reactions can either occur at the surface of the electrode or at the interface of the SEI with the bulk electrolyte. The composition of the reactive species in the electrolyte can vary, with the cyclic EC being the most-used organic solvent molecule along with the linear diethyl carbonate (DEC).7,48,49 It has been found that the SEI tends to be stable and act favorably in an EC excess.7,42,46 The true nature and structure of the SEI and its formation mechanisms are yet to be understood, and given its importance, it deserves more attention from the computational electrochemistry community. Electrochemical reactions that form the SEI are largely analogous to reactions of importance in electrocatalysis, for example, CO2 reduction,50 where the potential dependence of elementary reactions is added to electronic structure calculations via the “computational hydrogen electrode” technique.51 This technique relates the chemical potential of a proton/electron pair, at 0 VRHE, to that of H2 gas, that is, through the reference electrode reaction. Here, we employ an analogous computational lithium electrode technique, relating a Li+/e− pair to that of bulk Li via the reference electrode reaction Li+ + e− ↔ Li(s) that defines 0 VLi/Li+. This allows us to study the electrochemical steps of SEI formation while incorporating potential in a thermodynamically consistent manner, allowing for static electronic structure calculations of reactions and, ultimately, the creation of potential-dependent free energy diagrams. We employ this technique (within a density functional theory framework) to explore the mechanisms and driving forces of SEI formation. We study the lithium metal surface to avoid the ambiguous structural terminations and chiralities possible in graphitic systems; therefore, this study applies primarily to lithium-metal batteries, but we expect the general conclusions to also hold for Li-ion batteries. We also study reactions on Li2O and Li2CO3 surfaces, which are two of the most commonly reported inorganic Li compounds in the SEI structure. We demonstrate that the CLE technique can correctly predict the formation of the anode inorganic−organic Li structure and provide insights into the working mechanisms of the potential-dependent reactions, explaining the commonly accepted image of the SEI structure.7,27,43,52−55 We expect that the insights provided by this study will rationalize experimental observations and set the stage for future research aimed at tailoring the SEI structure through mechanism-based control

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METHODS Computational Lithium Electrode Technique. Consider the following reaction of ethylene carbonate (EC) with a solvated lithium ion (Li+) and an electron (e−) at an electrode surface (*) to form an adsorbed complex (Li−EC−*): EC + Li+ + e− → Li − EC − *

(1)

The change in free energy of this reaction is the difference between the chemical potentials of the reacting species in each side. While calculating the chemical potentials of the uncharged species (EC, *, and Li−EC−*) from electronic structure calculations is relatively straightforward, the calculation of the solvated lithium ion and the “extra” electron is less so. To this end, we use a direct extension of a standard methodology in electrocatalysis to calculate the combined chemical potential of this pair. In the past decade, electronic structure calculations have been successfully used in the field of catalysis to calculate the reactivity of small molecules and organic species on electrode (electrocatalyst) surfaces.50,51,56,57 A significant breakthrough in the field was made by Nørskov et al. who suggested a concise way to thermodynamically account for the reactive potential of the surface.51 This method is known as the “Computational Hydrogen Electrode” (CHE), and it assumes that, at zero volts versus the reversible hydrogen electrode, the reactive proton−electron pairs (H+ + e−) and gas-phase H2 molecules are defined to be in equilibrium (reaction 2), and therefore one can use their chemical potential interchangeably.50,51 + 2H(aq) + 2e− F H 2(g)

(2)

In this work, we use the same framework but reference the energies to the standard lithium electrode, which is defined by the equilibrium between bulk lithium metal and Li+/e− pairs. This computational lithium technique58 assumes that, in the defined reference electrode, the Li/Li+ pair is in equilibrium, and therefore, analogous to reaction 2, we can define the following: Li+ + e− F Li(bulk)

(3)

where Li(bulk) refers to the bulk lithium in the body-centered cubic (bcc) structure. Thus, we simply replace the chemical potential of the Li+ and e− (at 0 V vs Li/Li+, VLi/Li+) by that of bulk lithium. Analogous to electrocatalysis studies, this potential can be adjusted within the standard relation between free energy and electrical potential (G = −eU). As a result, we can map out competing reaction pathways and incorporate electrode potential dependence to clarify the conditions that favor and disfavor stable SEI formation reactions. In short, using this CLE technique, we can rewrite the free energy difference (ΔG) of an electrochemical reaction, such as reaction 1 at 0 V versus Li/Li+ as ΔG = μ[Li − EC − *] − μ[EC] − μ[*] − μ[Li+ + e−] = μ[Li − EC − *] − μ[EC] − μ[*] − μ[Li(bulk)] (4) B

DOI: 10.1021/acs.jpcc.9b03886 J. Phys. Chem. C XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry C If one desires to know the free energy diagram at a different potential, the values need only be adjusted by −eU for each Li+/e− pair. Atomic Structures and Electronic Structure Calculations. Electronic structure calculations were performed using density functional theory (DFT) as implemented in the Dacapo calculator59 within the Atomic Simulation Environment (ASE).60,61 Ultrasoft pseudopotentials62 and revised Perdew−Burke−Ernzerhof (RPBE) exchange-correlation functional63 were employed. Calculations were carried out on three representative surfaces: the Li bcc (100) surface using 3 × 3 × 3 unit cell (27 atoms), the Li2O fcc (111) surface using 2 × 1 × 2 unit cell (48 atoms; Li32O16), and the Li2CO3 monoclinic (010) surface using 2 × 1 × 2 unit cell (96 atoms; Li32C16O48). Periodic boundary conditions were applied in the x and y directions, while a 15 Å vacuum was added in the z direction, orthogonal to the surfaces. To avoid molecule− molecule interactions for the longest adsorbed molecule (LiEDC) on Li bcc (100), we used a larger unit cell for this surface (4 × 4 × 3 unit cell; 48 atoms). A Monkhorst-Pack kpoint sampling64 of (3,3,1) in the (x,y,z) directions was used for the Li (100) and the Li2O (111) surfaces, while a (2 × 2 × 1) sampling was used for Li2CO3 (010). The values for the planewave and density cutoff energies were adopted as 340 and 500 eV, respectively. For geometry optimizations, only the top layer of the Li (100) slab and the adsorbate were free to relax, and the bottom 2/3 of the structure was kept frozen in order to mimic the bulk structure. We verified that the results did not meaningfully change while fixing only the bottom 1/3 of the surfaces. In the case of the Li2CO3 (010) surface, however, we kept the bottom half of the surface frozen due to its complex structure. The structures were considered fully relaxed when the magnitude of the individual forces on all unconstrained atoms were lower than 0.05 eV/Å. Free energies of the adsorbates were calculated through normal-mode analysis assuming only the vibrations of the adsorbed species (adopting the harmonic oscillator approximation) and ignoring entirely those of the surfaces.61,65 We disregarded solvation effects in this study as these are challening to include rigorously; however, this is equivalent to an assumption that solvent corrections will be similar and will, to a large degree, cancel out. Of course, future studies may include such effects and other solution-side effects, such as salt concentrations. The free energies of the free molecules were calculated based on the ideal-gas approximation using the experimental/realistic vapor pressures for the species. Specifically, the vapor pressure used for the EC molecule was 0.0098 mmHg,66 while that for the ethylene molecule was 52,100 mmHg.67

where * indicates the reactive surface or a molecule chemisorbed to such. Through such reactions, the EC molecule can be decomposed (in combination with lithium) to common organic (LiEDC) and inorganic (Li2CO3 and Li2O) products. In the following, we will comprehensively study the different pathways for EC degradation, including the effect of potential and the surface structure on the free energy profiles of each pathway and the possibility of related nonelectrochemical steps. Li Metal Surface. We start with a discussion of mechanisms on the bcc (100) surface of Li, chosen since previous studies have shown that tetragonal Li (100) has the lowest surface energy compared to other structures such as orthorhombic Li (110) and Li (111).68−70 On this surface, we studied different intermediate pathways that occur with EC, first a purely electrochemical mechanism that consumes Li+/ e−, and second, a non-electrochemical/electrochemical (mixed) mechanism that is possible due to the availability of metallic lithium in this system. Figure 1 provides a schematic of these two classes of reaction mechanisms. In the electrochemical mechanisms (part

RESULTS AND DISCUSSION Ethylene carbonate is perhaps the most prevalent battery electrolyte compound, and here we will study its decomposition products at a negative lithium electrode. Experimentally, the SEI tends to build up while the battery is charging; that is, when the negative electrode (commonly referred to as the anode) acts as a cathode. With current flowing in this direction, elementary reduction reactions such as the following can be expected to take place,

Figure 1. Schematic representation of the two main reaction mechanisms for LiEDC formation from EC: (a) electrochemical mechanisms and (b) mixed mechanism.

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C2H4CO3 + Li+ + e− + * → LiCO3C2H4*

(5)

C2H4CO3 + Li+ + e− + * → LiCO3* + C2H4

(6)

(a)), an EC molecule is reduced by an electron from the Li surface (the negative electrode), combining with a Li+ to preserve charge neutrality and results in either an adsorbed open-ring Li-EC structure (Li-EC-*) or simply carbonate precipitation (Li-CO3-*) accompanied by liberation of ethylene gas. (Note that both ethylene gas and Li2CO3 are commonly observed under operating conditions in lithiumbased batteries.) In a second step, another EC molecule along with a Li+/e− pair can result in the formation of adsorbed LiEDC (possibly with the generation of a C2H4 gas molecule, C

DOI: 10.1021/acs.jpcc.9b03886 J. Phys. Chem. C XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry C depending on the first step). These are fully electrochemical mechanisms, that is, each elementary step involves an electron transfer from the circuit, that directly follow the forms such as reactions 5 and 6. Thus, the driving force for each of these steps will have a direct response to the electrode potential. Additionally, there are experimental evidences that the SEI starts forming immediately after the electrode is put in contact with the electrolyte whether under the open circuit voltage71,72 or even by merely immersing it into the electrolyte solution.13 Thermodynamically, this is not surprising as the solid lithium metal has equivalent reducing power to a Li+ + e− pair at 0 VLi/Li+. Figure 1b suggests a mixed mechanism in which an EC molecule interacts directly with the Li (100) surface and results in carbonate precipitation and ethylene liberation through a non-electrochemical reaction. In a subsequent step, an EC molecule along with a Li+/e− pair will react on the CO3-covered Li surface, resulting in an adsorbed LiEDC structure. Thus, based on the proposed mixed mechanism, the LiEDC can be formed in a single electrochemical step on the Li metal surface. It is worth mentioning that, in principle, both steps could be non-electrochemical. However, as we will show shortly, the first step has a strong tendency to continue until it fully carpets the Li surface with EC molecules, denying access to surface Li atoms in order to complete the LiEDC formation. On the other hand, we know that SEI formation is accelerated during charging cycles, which indicates that reactions with at least one electrochemical step are dominant. Therefore, we will consider the second step to be electrochemical using the electrolyte Li+ ions. To study these proposed mechanisms, we use the aforementioned CLE technique for the EC molecule’s electroreduction by examining all reasonable structures on the Li (100) surface at 0 VLi/Li+. These calculations show that breaking EC into a precipitated carbonate group and ethylene gas is much more favorable than a direct precipitation into LiEC-*. Our calculations show that the EC precipitation is slightly more favorable through the non-electrochemical step (−2.62 eV vs −2.46 eV at 0 VLi/Li+); however, since their values are relatively close, we will present our results for both steps to provide a more comprehensive overview of the subject. First, we analyze the free energy diagram for the electrochemical mechanisms, as it is depicted in Figure 2 along with the schematic representations of the calculated structures. We found that the open-ring Li-EC structure on the Li slab (Li-

EC-*) results in an uphill free energy by 0.28 eV, which can be interpreted as an applied potential of about least −0.28 VLi/Li+ would be necessary for this step to occur, whereas experimentally, the SEI is known to form at potentials positive of 0 VLi/Li+.7 Meanwhile, the competing mechanism, that is, carbonate adsorption (forming Li-CO3-*) and ethylene liberation, is strongly exergonic, being downhill in free energy by more than 2 eV at 0 VLi/Li+. Based on this strong difference and allowing for only the electrochemical mechanisms, the first EC molecule on the Li (100) surface will most probably proceed through the carbonate precipitation/ethylene liberation step. Adding a second EC on the Li-CO3-* structure can result in the formation of an LiEDC molecule either in the vertical (LiEDCv-*) or parallel (LiEDCp-*) orientation with respect to the surface plane, with ΔG = 0.88 and 0.66 eV, respectively. Note that ignoring solvation effects may obscure this small difference in binding energies. At the end, although LiEDC formation would be, in net, downhill in the free energy, this uphill step would thermodynamically preclude this pathway as a dominant one. The only reaction that exhibits negative formation free energy after the addition of a second EC molecule is a subsequent CO3 adsorption (2(Li-CO3)-*) with C2H4 gas liberation (ΔG = −2.45 eV). Therefore, we can argue that the LiEDC formation through the electrochemical mechanisms will not be directly favorable on Li metal surfaces and carbonate precipitation will dominate. For the mixed pathway, we observed the same general trend. Namely, after the first carbonate precipitation through a nonelectrochemical reaction (ΔG = −2.62 eV), the LiEDC formation was uphill through both electrochemical and nonelectrochemical mechanisms by ΔG = 1.09 and 0.45 eV, respectively, while the second carbonate precipitation has ΔG = −2.90 eV. A natural next step is to explore the extent to which CO3 deposition remains favorable and whether LiEDC forms favorably on a CO3-saturated surface. To examine this, we chose the non-electrochemical reaction, which was slightly more favorable, having in mind that the general mechanism is the same for both cases. Figure 3 clearly shows that carbonate precipitation continues to be strongly exergonic until the surface is saturated with carbonate, which, in this system, is reached with four carbonates per nine lithium surface atoms.

Figure 2. Free energy diagram of the first and second EC molecule on the Li (100) surface at 0 V versus Li/Li+ according to the electrochemical mechanisms of Figure 1. The values denote the free energy difference (ΔG) between the steps. Purple, gray, red, and white spheres correspond to Li, C, O, and H atoms, respectively.

Figure 3. Free energy diagram of EC molecules on Li (100) surface at 0 V vs Li/Li+. The values denote the free energy difference (ΔG) between the steps. Purple, gray, red, and white spheres correspond to Li, C, O, and H atoms, respectively. D

DOI: 10.1021/acs.jpcc.9b03886 J. Phys. Chem. C XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry C Even on this fully CO3-saturated Li surface, LiEDC formation remains unfavorable. Interestingly, carbonate precipitation is still favorable even for a fifth EC on this surface. Within the constraints of our unit cell, at this stage the lithium-carbonate layer begins to exfoliate, suggesting the favorability of forming a separate Li2CO3 phase. (To unambiguously identify whether this occurs via an exfoliation mechanism would require much larger unit cells; however, we saw the same effect when we doubled the unit cell.) This suggests that the material is being driven toward Li2CO3, and we can readily calculate the driving force for this from EC and various lithium precursors with the computational lithium electrode technique, that is | l 2Li(s) o o o o o o o o o o o + − o Li (Li e ) + + EC + m → Li 2CO3(s) + C2H4 } (s) o o o o o o o o o o 2(Li+ + e−) o o n ~

Figure 4. Free energy of carbonate addition on the Li (100), Li2O (111), and Li2CO3 (010) surfaces at 0 V versus Li/Li+ through nonelectrochemical (black line) and electrochemical (red line) mechanisms. Purple, gray, red, and white spheres correspond to Li, C, O, and H atoms, respectively.

which we find to be strongly favorable with ΔG = −3.26 eV at 0 VLi/Li+. These calculations suggest that LiEDC formation directly on lithium metal surfaces is not thermodynamically favorable but that the continuous deposition of carbonate should result in the formation of inorganic lithium compounds. This helps to explain the common experimental observations of compounds such as Li2CO3 and LiO2 as inner SEI products. Inorganic Li Surfaces. As discussed in the previous section, the most favorable reaction on metallic lithium surfaces is the consecutive precipitation of carbonate groups, which we anticipate would favorably lead to the formation of an inorganic Li2CO3 phase with an overall ΔG of −3.26 eV. Interestingly, experimental studies have shown that inorganic Li structures like Li2CO3 and Li2O can be detected in the SEI on both lithium-metal and lithium-ion electrodes,7,42,49 and theoretical works have also suggested the possibility of Li2CO3 formation through several EC reactions in Li-rich environments.47 We estimate the formation energy of Li2O from 2 Li+ + 2 e− + EC (liberating C2H4 and CO2) to be −2.36 eV at 0 VLi/Li+, again calculated with the CLE technique. This logically leads us to examine further reactions on both inorganic Li surfaces as alternative locations for LiEDC formation. Although these SEI compounds are insulators by nature, experimental impedance studies show that they do not immediately stop the electronic transfer but rather slowly decrease it upon building more layers.71,73 Consequently, those surfaces can still facilitate electrochemical reactions although with a constantly decreasing rate. Therefore, in the following, we study the plausibility of LiEDC formation on the Li2CO3 (010) and Li2O (111) inorganic surfaces. (The Li2O (111) surface is the most thermodynamically stable one,74−77 while the Li2CO3 (010) surface is the one with the highest Li+ diffusivity.78,79) We first examine the precipitation of carbonate (from EC, with concomitant C2H4 liberation) on these surfaces through both electrochemical and non-electrochemical mechanisms, shown in Figure 4. As discussed earlier, the non-electrochemical step is energetically preferred for the case of the Li surface (ΔG = −2.62 eV vs −2.46 eV). However, the electrochemical step became more favorable for the Li2CO3 (010) and Li2O (111) surfaces. We can see a correlation between the percentage of the surface Li atoms and favorability of the two routes’ electrochemical reaction, as shown in Figure 4. In the materials

with the lowest lithium percentage, the electrochemical step seems to be favored, while with higher Li percentages, the opposite is true. This can be easily understood: materials such as Li2CO3 has a low lithium content and benefits from the addition of extra Li with the carbonate as in the electrochemical mechanism. The lithium metal already has a surplus of lithium, and thus extra lithium deposition is not needed to form lithium carbonate. More precisely, the reaction free energy of the carbonate on the Li2O (111) surface (which consists of 67% Li atoms) is −2.48 eV for the nonelectrochemical step and −2.62 eV for the electrochemical one. For the Li2CO3 (010) (consists of 33% Li atoms) surface, the carbonate precipitations exhibit ΔG = 0.81 and −1.40 eV for the non-electrochemical and electrochemical reactions, respectively. Therefore, the non-electrochemical mechanism is favored on the Li-rich surfaces like Li (100), while the electrochemical mechanism is preferred as we move to lower Li percentage on the surfaces. Nevertheless, based on these calculations, we only consider eletrochemical mechanisms for the EC reduction on our inorganic surfaces. For convenience, Table 1 represents all the reactions we considered for the first and second EC reduction on the Li2CO3 (010) and Li2O (111) surfaces. The free energy diagram of these reactions on the Li2CO3 (010) surface at 0 Table 1. Electrochemical Reactions of the First and Second EC Molecule (First and Second Electron, Respectively) on Inorganic Li Surfacesa

a

E

* is the Li2O (111) or Li2CO3 (010) surfaces. DOI: 10.1021/acs.jpcc.9b03886 J. Phys. Chem. C XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry C VLi/Li+ is shown in Figure 5. We can see that the first EC molecule reacting on the Li2CO3 (111) surface can result in

0.28 eV. We should note here that the lithium deposition on the Li (100) surface results in a higher free energy (0.54 eV). Interestingly, electrodepositing the open-ring Li-EC structure on the slab (Li-EC-*) exhibits a downhill free energy pathway of −0.61 eV; this pathway was always uphill on the lithium surfaces described earlier. However, the most favorable elementary step for the first EC electroreduction is again carbonate precipitation (Li-CO3-*), liberating an ethylene gas molecule (ΔG = −2.62 eV). . Upon the addition of the second EC molecule, we can again observe a fully downhill path for LiEDC-* formation through the Li-EC-* formation in the first step and Li-CO 3 precipitation in the second (ΔG = −1.68 eV). Finally, the two most stable formations are Li2-CO3-* and 2(Li-CO3)-*. The former one is the product of either the Li-* structure along with an ethylene gas molecule (ΔG = −2.95 eV) or the Li-CO3-* structure with an EC liberation (ΔG = −0.05 eV). The latter one is formed only as a subsequent step from the LiCO3-*, liberating a C2H4 gas molecule (ΔG = −3.17 eV). Thus, the LiEDC structure and the LiBDC structure can be formed on the Li2O (111) surface through two downhill electrochemical steps, with LiEDC exhibiting a slightly lower free energy compared to the LIBDC. However, the structures with the most downhill electrochemical pathways are the carbonates precipitated and especially the 2(Li-CO3)-*, indicating that the lithium carbonate structures are more likely to be formed on the Li2O surface. Experimental80−83 and theoretical studies84−86 have also mentioned that structures like Li2CO3, Li2O, LiEDC, C2H4, and CO3 are found in the SEI of Li-based batteries.

Figure 5. Free energy diagram of the first and second EC molecule on the Li2CO3 (010) surface at 0 V versus Li/Li+ according to the electrochemical reactions in Table 1. The values denote the free energy difference (ΔG) between the steps. Purple, gray, red, and white spheres correspond to Li, C, O, and H atoms, respectively.

three structures: the Li-* with ΔG = 0.77 eV, the Li-EC-* with ΔG = 0.18 eV, and the Li-CO3-* with ΔG = −1.40 eV. Therefore, carbonate precipitation is the most favorable reaction on this surface as well. Reducing a second EC molecule on these three structures can result in five different cases (the Li2-EC is unstable on this specific surface). Here, we finally observe a fully downhill path for LiEDC-* formation through Li-CO3-* formation in the first step and Li and EC precipitation in the second (ΔG = −2.02 eV). The second EC molecule on the Li-* structure gives as products the Li2-* (ΔG = 0.45 eV) or Li2CO3-* (ΔG = −4.12 eV) structures. Also, for the case of the Li-EC-* structure, after the additional EC molecule, we found that the LiEDC and the LiBDC ((C2H4OCO2Li)2) structures are formed with free energies equal to −3.60 and −3.78 eV, respectively. Therefore, the LiEDC structure can be formed on the Li2CO3 (010) surface through two downhill electrochemical steps. Similarly, we have studied the first and second EC molecules’ electroreduction on the Li2O (111) surface at 0 VLi/Li+ (Figure 6). Expectedly, the adsorption of Li atom on the Li2O (111) (Li-*) surface results in an uphill free energy of

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CONCLUSIONS In this work, we demonstrated that the computational lithium electrode (CLE) technique is a powerful tool to study electrochemical reactions leading to the solid electrolyte interphase (SEI) formation on the batteries’ negative electrode. In this approach, we replace the free energy of the Li+/e− pair with that of bulk Li under the assumption that they are in equilibrium at the reference electrode, which enable us to create potential-dependent free energy diagrams. Using CLE, we showed that the decomposition of ethylene carbonate into ethylene gas and adsorbed carbonate is a favorable reaction under many conditions, suggesting a route to the often-observed products of Li2CO3 and ethylene gas. This reaction can be triggered either electrochemically or nonelectrochemically, that is, reacting directly with excess lithium in the surface. We did not see direct routes to the organic precursor LiEDC on pure lithium metal surfaces. When we examine analogous reactions on the inorganic Li2CO3 and Li2O surfaces, we again see carbonate precipitation (associated with ethylene liberation) as a favorable reaction mechanism. However, under these conditions, we did observe favorable pathways to organic compounds like LiEDC and LiBDC. These results computationally reaffirms the known anode−inorganic−organic Li compound picture of the SEI structure and establishes the CLE as a useful method in this context. Moreover, these calculations shed insight into the reaction pathways responsible for solvent degradation and SEI formation in lithium-based electrodes. We expect that similar conclusions will apply on graphitic systems. This study opens the door toward the exploration of many other effects, such as

Figure 6. Free energy diagram of the first and second EC molecule on the Li2O (111) surface at 0 V versus Li/Li+ according to the electrochemical reactions in Table 1. The values denote the free energy difference (ΔG) between the steps. Purple, gray, red, and white spheres correspond to Li, C, O, and H atoms, respectively. F

DOI: 10.1021/acs.jpcc.9b03886 J. Phys. Chem. C XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry C

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the composition of the electrolyte and the inclusion of reaction barriers in such reactions. We expect such studies to help understand why some organic molecules work better than others to form stable SEI’s and accelerate the pace of materials discovery in this essential field of energy storage.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. ORCID

Andrew A. Peterson: 0000-0003-2855-9482 Author Contributions ‡

M.A.G. and J.H. contributed equally to this work.

Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This material is based on the work supported by the U.S. Department of Energy through the DOE EPSCoR program under award no. DE-SC0007074.

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REFERENCES

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DOI: 10.1021/acs.jpcc.9b03886 J. Phys. Chem. C XXXX, XXX, XXX−XXX