Degradation in lithium ion battery current collectors

The release of electrons into an external circuit results from the pair of chemical 'half-reactions' that occur at the electrodes. These reactions are in principle reversible, giving the cell its rechargeable nature. The difference in voltage between the couples of both electrodes (for example Fe²⁺/Fe³⁺ paired with Li/Li⁺ in an LiFePO₄/Li cell) determines the overall cell voltage. The energy density of the cell can be calculated according to equations (1) and (2) for the gravimetric and volumetric energy densities respectively, where $\bar V$ represents the average voltage and Q represents charge [9]:

$$\begin{equation}{E_{{\text{grav}}}} = \frac{{\bar V \times Q}}{{{\text{mass}}}}\end{equation} \tag{ 1 }$$

$$\begin{equation}{E_{{\text{vol}}}} = \frac{{\bar V \times Q}}{{{\text{volume}}}}\end{equation} \tag{ 2 }$$

$$\begin{equation}\bar V = \frac{{{\text{ }}\int\limits_0^Q {V{\text{d}}Q} }}{Q}.\end{equation} \tag{ 3 }$$

Most commercial cells are composed of an LiCoO₂ cathode and graphite anode [36]. State-of-the-art EVs employ LiNiMnCoO₂ or LiNiCoAlO₂ cathodes and a lot of research is focused on improving the performance of these materials [37]. The electrode active materials are mixed with a high surface area carbon black (to improve conductivity) and a polymeric binder, generally polyvinylidene fluoride at the cathode (to maintain contact between all components) and coated onto the current collector. The cell is assembled in a discharged state, with all Li⁺ ions in the host oxide at the cathode. Upon application of a potential difference, the Li⁺ ions deintercalate from the transition metal oxide and are inserted into the graphite at the anode. Upon discharge, Li⁺ ions are removed from the graphite and intercalate back into the metal oxide host framework. These intercalation reactions result in the production and consumption of electrons which move through an external circuit to balance the charge. The anodic, cathodic and overall cell reactions are as follows [9]:

$$\begin{equation}{\text{Anodic Reaction}}:{\text{L}}{{\text{i}}_x}{\text{C}} \to {\text{C}} + x{\text{L}}{{\text{i}}^ + } + x{{\text{e}}^ - }\end{equation} \tag{ 4 }$$

$$\begin{equation}{\text{Cathodic Reaction}}:x{\text{L}}{{\text{i}}^ + } + x{{\text{e}}^ - } + {\text{L}}{{\text{i}}_{1 - x}}{\text{Co}}{{\text{O}}_2} \to {\text{LiCo}}{{\text{O}}_2}\end{equation} \tag{ 5 }$$

$$\begin{equation}{\text{Overall Reaction}}:{\text{L}}{{\text{i}}_x}{\text{C}} + {\text{L}}{{\text{i}}_{1 - x}}{\text{Co}}{{\text{O}}_2} \to {\text{C}} + {\text{LiCo}}{{\text{O}}_2}.\end{equation} \tag{ 6 }$$

Intercalation type cathodes have dominated the commercial LIB market since their first commercial implementation by Sony [38]. Such cathodes have a host structure that enables the reversible hopping of Li⁺ ions for example 1D tunnels in the case of LiFePO₄ or 2D layers in the case of LiCoO₂ [9]. Such materials are of interest as they are able to maintain their structure upon removal and insertion of lithium, allowing for good capacity retention as a function of cycling.

Lithium metal could be used as the anode in an ideal case as it provides a high specific capacity of 3680 mAh g⁻¹ [39]. Unfortunately, lithium metal does not plate and strip uniformly as a function of cycling and it forms dendrites at the electrode surface—as these grow, they may pierce the separator, coming into contact with the cathode and causing a potentially dangerous short circuit [39]. Graphite is therefore mostly employed at the anode as it avoids this safety concern, although it also provides a lower energy density. It should be noted however, that lithium dendrite growth has still been observed on graphite [40].

LIB electrolytes consist of a lithium salt (generally LiPF₆) in a mixture of organic carbonate solvents. These solvents are usually some combination of ethylene carbonate (EC), ethyl methyl carbonate (EMC), diethyl carbonate (DEC) or dimethyl carbonate (DMC). These carbonates have low viscosity, poor electronic conductivity, good ionic conductivity and have relatively good chemical stability [39]. LIB cells are assembled in a dry environment (e.g. a glovebox or dry room) to avoid LiPF₆ salt hydrolysis, which forms hydrofluoric acid (HF). Cells can be assembled in a variety of configurations, with the most common experimental configurations being pouch cells and coin cells. Coin cells seal the cell components in a stainless-steel casing and pouch cells are polymer pouches that contain the components and have conductive foil tabs welded to the electrodes.

The rate at which a battery is charged/discharged is given in terms of the C-rate. A charge or discharge current of 1 C is the current required to charge or discharge an electrode in 1 h (assuming 100% theoretical capacity). A C-rate of C/2 and 2 C will charge or discharge fully in 2 and 0.5 h respectively.

An LIB cell's open circuit voltage (V_OC) is dependent on the electrode potentials of the anode and cathode, μ_A and μ_C respectively according to equation (7) [9]. Stable operation of the battery requires that this voltage lies within the stability window for the electrolyte, E_g, defined by its highest energy occupied molecular orbital (HOMO) and lowest energy unoccupied molecular orbital (LUMO) as described in equation (8) and seen in figure 2:

$$\begin{equation}{V_{{\text{OC}}}} = \frac{{{\mu _{\text{A}}} - {\mu _{\text{C}}}}}{{\text{e}}}\end{equation} \tag{ 7 }$$

$$\begin{equation}{E_{\text{g}}} = {E_{{\text{LUMO}}}} - {E_{{\text{HOMO}}}}.\end{equation} \tag{ 8 }$$

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If the LUMO of the electrolyte is lower than μ_A (the anode potential), it will be electrochemically reduced. Similarly, if the electrolyte HOMO is greater than μ_C (the cathode potential), it will be oxidised. In an aqueous electrolyte, the above reactions are hydrogen and oxygen evolution respectively, limiting the maximum voltage of an aqueous cell to less than 2 V [36]. Commercial batteries therefore employ non-aqueous electrolytes which have a wider E_g in order to maximise the cell voltage and hence energy density.

The μ_A of a lithium metal or graphite anode lies beneath the LUMO of most non-aqueous organic electrolytes [39]. The electrode's operation is therefore dependent on the formation of a 'passivation' layer, known as the SEI [41]. Upon initial application of a potential, the electrolyte is reduced, forming the SEI, an electronically insulating but ionically conductive layer. These properties prevent further reaction of the electrolyte with the anode, allowing the cell to operate at voltages greater than 4.0 V. Similar layers may also form on LIB current collectors, influencing their corrosion behaviour.

Reactive metals form a thin passive layer on their surfaces under ambient or submerged conditions—typically oxides or oxy-hydroxides. They kinetically restrict dissolution and further reaction of metals and alloys despite the strong thermodynamic driving force for oxidation. These films can vary in thickness, composition and properties, depending on the metal on which they form. Semi-conducting passive films form on Fe, Cr, Ni and Cu and will only grow up to a few nanometres in thickness. Insulating oxide films grow on so-called valve metals such as Al, Nb, Ti and Ta and these materials can be further oxidised by application of an external voltage (anodisation) to produce significantly thicker surface films. Passive films exhibit a potential drop at both the metal–film and film–electrolyte interfaces, as well as within the passive layer. The electric field strength within the passive layer is therefore extremely high and enables the migration of ions through the film at room temperature [43]. Although their properties and formation are different, it is interesting to note the similarity between the SEI on a battery electrode and a passive film on a reactive metal.

The structure and chemistry of passive films have been extensively studied and they have been shown to have complicated structures made up of nanoscale layers of oxides and hydroxides. The structures and compositions are dependent on electrode potential (and redox history), metal and electrolyte composition. The native/passive oxide on copper is usually proposed to be crystalline and a combination of copper I and II oxides. The native/passive oxide that forms on aluminium is very temperature dependent and can even be amorphous [43].

Pitting corrosion is a type of localised corrosion that takes place at passivated metal surfaces in the presence of 'aggressive' anions, such as halides [43]. Pit formation generally happens in two phases and is depicted in figure 3: breakdown of passivity and pit initiation/pit growth—the growth being particularly stabilised by halide species. Pit growth stops upon repassivation. Pitting is generally considered to be an autocatalytic process: once a pit starts to grow, local aggressive conditions develop such that further pit growth is accelerated. The dissolution remains localised and the anodic and cathodic reactions separate spatially on the metal surface during corrosion.

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There are various theories to describe the initial breakdown of passivity—the initiation phase—and these may depend on the environmental conditions of the film. For metals, the initiation mechanisms can be categorised as penetration, film breaking or adsorption. For alloys, pits usually initiate a local heterogeneities, such as inclusions, secondary phases or mechanical damage [44]. It is the least understood aspect of pitting and there is no consensus of a generalised mechanism. The critical factors affecting pitting corrosion in aqueous systems have been thoroughly reviewed by Frankel [44].

Crevice corrosion is another form of localised corrosion that can occur in specific geometries or as 'under-deposit'. It occurs when a part of a metallic surface is in contact with a small volume of confined, stagnant liquid but most of the surface is in contact with the bulk environment [45]. Such an area is found in a 'crevice' as seen in figure 4. Zones susceptible to crevice corrosion may arise from a variety of situations. These include a limited mass transport by diffusion or convection between the inside of the crevice and the bulk environment, and significant solution resistance between the inside and outside of the crevice.

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The change of the local environment inside a crevice zone into more aggressive conditions results in crevice corrosion. On passivated metals, these conditions must become severe enough to breakdown the passive film and begin metal dissolution. Such a local change in environment takes place in two steps.

Firstly, the crevice experiences a local accumulation of corrosion products or exhaustion of reactants due to restricted transport kinetics. A Galvanic cell is then developed between the crevice and the external surfaces. This in turn results in a difference in corrosion potential on the inside and outside of the crevice due to solution resistivity (known as IR drop, where I is current and R is resistance). The build-up of the Galvanic cell also further modifies the crevice environment due to electrolytic migration between the crevice and bulk solution. Once the local environment becomes severe enough to break passivity, dissolution damage occurs in the crevice.

The severity of the crevice corrosion is dependent on the crevice geometry, external surface composition, IR drop (which is dependent on crevice geometry and solution conductivity), electrolyte composition and temperature.

Copper foils used as current collectors have a native oxide layer, which would be formed in air. Figure 5 shows a cyclic voltammogram (CV) of a copper current collector foil undergoing an initial polarisation. The CV is measured in an LiPF₆ based electrolyte, at 40 °C and in the range of 0–5 V at a scan rate of 10 mV s⁻¹ [47]. The CV begins at open circuit potential (3.326 V) and is initially polarised in the cathodic direction. A peak is observed at ∼3.1 V which is thought to correspond to decomposition of HF impurities contained in the electrolyte [30]. Other reports indicate that HF decomposition takes place between ∼1.8 and 2.4 V via HF dissociation and subsequent hydrogen evolution and LiF deposition [50]. In this case, the peak observed at ∼3.1 V is more likely to correspond to copper reduction. A large cathodic peak is also observed in the range of 1.5–3 V. This peak is thought to be due to the reduction of the native copper oxide film to metallic copper and the reduction of the electrolyte solvent, resulting in SEI formation on the metal surface [30, 47, 51]. However, these effects have not been studied extensively and their elucidation would be valuable for understanding copper current collector degradation and its contribution to overall cell capacity fading. No anodic reactions are observed up to about 3.7 V. Further polarisation to higher potentials results in an abrupt increase in current due to copper dissolution into the electrolyte. Myung et al investigated the surface of the copper metal with time of flight secondary ion mass spectrometry (ToF-SIMS) after polarisation in an LiPF₆ electrolyte and found that no significant passivation layer had formed [47]. Note that the sample was transferred through air. More recently however, Menkin et al observed an SEI layer formed on copper current collectors both after exposure to an LIB electrolyte and electrochemical cycling [52]. The implications of these findings for LIBs are discussed in more detail in this section.

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Over-discharge is detrimental to the copper foil current collector on the anode. Under normal operation of an LIB cell, over-discharge should not occur. However, in many LIB applications, including EVs, the cells are connected in series, forming a battery pack, in order to provide the high amounts of energy and power required. A significant difference in cell capacity can be present within such a pack, broadening the so-called capacity band and limiting the battery's capacity to that of the lowest capacity cell. This means that during battery operation, variable voltages will be experienced by the various cell's electrodes, and some will even experience over-charge and discharge. Over-charge and discharge can also occur during faulty operation of an LIB [53 – 57].

As discussed previously, a cell's potential is calculated by the difference in potential of the anode and cathode. Normal operation of an LIB is dependent on the materials used at the anode and cathode, however is generally ∼2.5–4.2 V. This means that the potential experienced by the copper anode varies between 0 and 1.5 V. At such potentials, copper should be cathodically protected (thermodynamically below its oxidation potential) and not subject to dissolution. The open circuit voltage of an LIB is generally above 3 V [26, 58]. The production of univalent copper takes place as follows in a non-aqueous environment:

$$\begin{equation}{\text{Cu}} \to {\text{C}}{{\text{u}}^ + } + {{\text{e}}^ - }.\end{equation} \tag{ 12 }$$

Equation (12) has a thermodynamic equilibrium potential 3.566 V in aqueous electrolytes under standard conditions [18]. Under non-aqueous conditions, this thermodynamic data no longer holds [18]. The potential at which equation (12) takes place has been measured experimentally and is listed in table 1. In one case it is measured to be at around 3.92 V in an LIB electrolyte, well away from the potentials experienced by the anodic current collector during standard LIB operation [54]. Copper dissolution is therefore mostly observed in LIBs upon over-discharge, when the anode potential reaches voltages greater than ∼3.5 V [18, 47, 54, 58 – 62].

Figure 6 shows the potential of the anode, cathode and cell during (a) normal operation and (b) upon 20% over discharge (120% depth of discharge) of an LiFePO₄ vs. graphite cell, with a lithium strip reference electrode. During normal operation, the anode reaches potentials between 0.2 and 0.9 V. Upon over-discharge, the anode potential increases sharply to more than 4 V. At these potentials, copper oxidation/dissolution is very likely to take place.

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Figure 7 shows a CV of copper in an LiPF₆ based electrolyte, with a platinum counter electrode and lithium reference electrode, measured over the potential range 2.50 V to 4.50 V. The oxidation peaks for Cu to Cu(I) and Cu(II) are observed at 3.92 V and 4.17 V respectively. The reverse reduction reactions can also be observed at 3.28 V and 3.19 V.

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This is just one example of copper oxidation and reduction in LIB electrolytes. Table 1 shows a list of copper oxidation and reduction potentials that have been measured for a variety of LiPF₆ based electrolyte systems.

Thermodynamically, copper dissolution is only expected at potentials greater than 3.5 V. In practice, the presence of water contamination can shift the threshold for copper oxidation to more negative potentials. The presence of impurities can lead to significant copper corrosion at OCP (discussed in further detail in section 5.2). Therefore, in an industrial setting, lithium ion cells' formation protocols include what is known as a tap charge. During a tap charge, cells are charged up to and held at a cell potential of 1.5 V in order to insert a small amount of lithium ions into the graphite on the anode, driving the anodic potential negative of the copper corrosion threshold [64]. This should be done as soon after electrolyte filling as is practical and before the cells are rested for electrolyte wetting. This tap charge step can be overlooked by academic battery researchers, however, it is an important step for minimising copper dissolution and further cell degradation.

Once Cu has dissolved, it will cause a cascade of chemical, electrochemical and electrical phenomena elsewhere in the cell. Cu⁺ ions can be redeposited on the anode as Cu metal, potentially forming dendrites that can grow and eventually penetrate through the separator, resulting in short circuit and cell failure [18]. Copper has correspondingly been detected throughout over-discharged LIB cells by various techniques. Zhang et al detected copper on the anode with energy dispersive x-ray spectroscopy (EDS) (as seen in figure 8) and x-ray powder diffraction (XRD) [65]. Li et al observed copper on both the anode and cathode with inductively coupled plasma—optimal emission spectrometry [28]. Maleki et al, using EDS, observed copper on both sides of the separator with EDS, confirming the proposal that copper is able to crossover through the separator [34]. Further similar observations have since been made of copper on both sides of the separator following over-charge, again however, the implications on long term cycling are not well understood [53, 54, 66].

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Such copper dissolution and redeposition can have significant consequences on the operation of the LIB. As mentioned already, the redeposition of metallic copper is thought to result in dendrite growth. He et al observed a sharp temperature increase to 120 °C after over-discharging to 120% depth of discharge (DoD) of an LiFePO₄ cell, indicating short circuit [54]. They proposed that during over-discharge, Cu⁺ cations are formed and diffuse through the separator to the cathode, where they are reduced back to metallic Cu and form dendrites, eventually leading to short circuit. This shorting behaviour can lead to unpredictable cell cycling and possibly catastrophic failure [32 – 35].

The copper may also redeposit on the anodic SEI, modifying its 'passivating' properties and impeding lithium mobility. Maleki reported an impedance increase of 250% after over discharging to 0 V, accompanied by a capacity loss of almost 65% [34]. The presence of copper throughout the cell may have many consequences including modifying the gas evolution and consumption (and hence swelling) behaviour in the cell [28, 34, 66, 67].

The extent of degradation of cell performance has been reported to be dependent on extent of over-discharge and C-rate [53, 54]. An LiFePO₄ vs. graphite cell functioned for 110 cycles at 1 C when over-discharged by 5%. When the same cell chemistry was over-discharged by 20%, it only survived two cycles [53]. Interestingly, over-discharged LiFePO₄ vs. graphite cells cycled at 0.1 C formed significant copper deposits on the cathode surface whereas cells cycled at 1 C only formed minor deposits that were more uniformly distributed. This effect occurs because the copper will experience the potential resulting in corrosion for significantly longer when cycled more slowly.

It is also interesting to note that researchers have observed copper current collector foil degradation, even under standard cell operating conditions. Dai et al studied LiCoO₂ vs. graphite cells and observed a degradation of the foil's mechanical properties upon standard cycling at 0.5 C, 0.8 C and 1 C for 200 cycles that became more severe with increasing C-rate (as can be seen in figure 9) [46]. The mechanical properties affected include elastic modulus, fracture strength, elongation and adhesion strength between the copper and the graphite coating. This study also observed the formation of pits on the Cu surface upon standard cycling, indicating localised rapid dissolution. The number and size of pits also increased with increasing C-rate. Although the effect of these pits on electrochemical performance is not explicitly discussed here, such pits can have significant implications on the electrode coating adhesion and therefore electrical contact. This loss of contact would result in decreased cell capacity and inhomogeneous current distribution that could lead to dendrite growth. Further studies that correlate the extent of pitting in copper current collectors to the extent of degradation in cell performance would be valuable.

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The mechanisms of how this copper dissolution takes place and influences the capacity fading, gas evolution, general stability and long term performance of LIBs is not fully understood, however it is clear that understanding these effects would be beneficial for the improvement of LIBs.

Copper dissolution has been observed directly upon exposure to the LiPF₆ based LIB electrolyte. The condition and composition of the electrolyte were found to influence the extent of this dissolution. Zhao et al used atomic absorption spectroscopy to observe that copper dissolution increased by up to several hundred ppm when a copper foil was exposed to an electrolyte that had been aged for six months versus a fresh electrolyte [48]. This is believed to be due to the higher levels of HF present in the aged electrolyte (from LiPF₆ salt decomposition). Copper dissolution was also observed upon cyclic voltammetry in the potential range 0–3.0 V by bulk electrolysis measurements. It was found that copper dissolution increased considerably in electrolytes doped with 500 ppm H₂O and 1000 ppm HF in both graphite coated and non-coated copper current collectors [58, 63].

The presence of HF in the electrolyte can contribute to copper corrosion. Dai et al reported significant pitting and even micro-cracks in copper current collectors after 720 h of immersion in an LiPF₆ based electrolyte (as seen in figure 10) [49]. This pitting was attributed to water impurities: related to the fact that in the presence of trace amounts of water, copper oxide is formed (equation (13)). Additionally, the LiPF₆ salt decomposes in the presence of water to form HF (equation (14)). Dissociation of this HF leads to a significant drop in pH and significant dissolution of copper oxide (equation (15)). The EDS spectrum in figure 10 shows that Cu, F, O and P elements are detected at the pitting sites. This suggests that corrosion products of copper fluorides and copper oxides are formed and the LiPF₆ salt decomposes. The authors do not investigate the effect of HF on metallic copper:

$$\begin{equation}{\text{Cu}} + x{{\text{H}}_2}{\text{O}} \to {\text{Cu}}{{\text{O}}_x} + 2x{{\text{H}}^ + } + 2x{{\text{e}}^ - }\end{equation} \tag{ 13 }$$

$$\begin{equation}{\text{LiP}}{{\text{F}}_6} + {{\text{H}}_2}{\text{O}} \to {\text{LiF}} + 2{\text{HF}} + {\text{P}}{{\text{F}}_3}{\text{O}}\end{equation} \tag{ 14 }$$

$$\begin{equation}{\text{CuO}} + 2{\text{HF}} \to {\text{Cu}}{{\text{F}}_2} + {{\text{H}}_2}{\text{O}}{\text{.}}\end{equation} \tag{ 15 }$$

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Shu et al stored copper foils in an LiPF₆ based battery electrolyte at room temperature for one month [21]. Upon removal, EDS and x-ray photoelectron spectroscopy (XPS) detected a surface film on the foil, composed of copper compounds, carbonate groups, alkyl groups, fluorides and phosphates. This film is likely to have formed from spontaneous decomposition of the electrolyte salt and solvents and trace impurities in the electrolyte.

Recently, Menkin et al investigated surface films formed on copper current collectors in the context of 'anode-free' LIBs, which are cells that employ a thin layer of lithium metal electroplated on copper as an anode [52]. The study found via XPS, nuclear magnetic resonance (NMR) and ToF-SIMS that a native interphase layer (native-SEI) is formed instantly upon copper immersion in an LiPF₆ based electrolyte. This layer was found to mostly be made of LiF however its nature was found to be very dependent on the copper interface composition and homogeneity. Upon electrochemical cycling, the surface film at the copper interface was observed to grow (electrochemical-SEI) and parasitic lithium plating was observed with in situ NMR. The morphology of the plated lithium was dependent on the homogeneity of the native-SEI. Interestingly, preferential plating was found, with XPS and ToF-SIMS, to be dependent on the distribution of the ionic conductivity of the SEI compounds rather than the electronic conductivity. This uneven plating and stripping can lead to dangerous dendritic short circuits and so this study highlights the important implications of understanding the formation and nature of the SEI on copper current collectors [68]. ToF-SIMS also detected copper oxides in the electrochemical SEI however this did not seem to result in preferential plating. The role and nature of dissolved copper deposited in the SEI is not well understood, however would be of significant interest for further study.

Although the interactions of the electrolyte with the copper current collector are now being investigated, further study into this chemistry would be valuable. In particular, understanding what implications the copper SEI has on LIB electrochemical performance and how/whether it plays a role in protecting against dissolution would be of interest.

The microstructure of the copper foils has also been shown to influence the corrosion behaviour of current collectors in LIBs. Kim et al heated copper current collector foils (in air) to temperatures ranging from 100 °C to 500 °C [69]. Upon heating above 200 °C, amounts of Cu₃O₄ and CuO measured by XRD increased considerably, with a concurrent increased surface roughness. These copper oxides formed nanoscale needle-like features on the surface as seen in figure 11. Upon cycling these bare, heat-treated Cu films against lithium (at a constant current of 0.05 mA for 100 cycles), non-negligible capacities were observed for the foils heated above 200 °C. This extra capacity is attributed to the electrochemical reactions of the copper oxides formed upon heat treatment, with lithium ions via conversion reactions [69].

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Pitting has been observed on copper current collectors (see figure 12) [25]. Copper foils were polarised in LiClO₄ based electrolytes (1 mA cm⁻² for 300 s) and examined with scanning electron microscopy (SEM). The pitting direction was further examined with cross-sectional SEM. It was found that the corrosion direction is dependent on the crystallographic orientation of the copper grain—so called 'etch pitting'. Further, solid corrosion products were not observed in the etch pits or on surface defects, indicating that this anodic polarisation of copper causes a typical activated dissolution reaction. The implications of this copper dissolution are poorly understood but could have significant effects, on the properties and growth of the anodic SEI for example. It is therefore of interest to further explore the effects of microstructure on copper dissolution behaviour.

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The effect of the graphite coating on the electrochemical stability of copper current collectors has been investigated [58, 63, 70 – 72]. Most reports suggest that the graphite coating provides protection against copper dissolution. This could be due to different dissolution kinetics as the graphite coating may act as a diffusion barrier, protecting the copper surface. Zhao et al observed this protective effect against copper oxidation in potentiostatic polarisation experiments [63]. Additionally, CV experiments showed that the graphite coating shifted the copper oxidation by 50 mV positive of the uncoated copper (figure 13) [63]. The positive shift decreases with increasing scan number. The CV for the coated copper also shows an increase in current with increasing scan number, whereas the uncoated copper shows reproducible behaviour with cycles. The authors concluded that this is due to the initially oxidised ions being redeposited back on top of the graphite, meaning that uncoated copper is present in subsequent cycles and the geometric surface area of copper in contact with the electrolyte varies. The implications on the electrochemical performance of the anode after this redeposition of copper are poorly understood, nonetheless, it is clear that there is a significant effect.

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At potentials near 0 V, Li–Al alloying occurs, preventing aluminium from being used as a current collector for an LIB anode [1]. At higher voltages however, aluminium is relatively stable. Protective passivation layers are formed on aluminium current collector surfaces in air and in LIB conditions [47]. In LIB conditions, aluminium forms a passive fluoride layer, unlike most other metals (including copper) [47]. Without this passive layer, aluminium would be unstable during LIB cycling and result in significant aluminium corrosion. As mentioned, this corrosion can result in significant performance loss in LIBs, however the contribution from current collector corrosion to total performance loss is not well understood. Typically, LIBs operate at voltages much higher than the standard equilibrium potential for Al/Al³⁺ (in aqueous solutions this is 1.38 V) [15]. If the battery is operating at a potential more negative than 4 V, an air formed oxide layer can sufficiently protect the aluminium metal from corrosion. Upon application of potentials greater than 4 V, the salt in the battery electrolyte (typically LiPF₆) breaks down and forms a layer of AlF₃, which further contributes to the corrosion resistance of aluminium [75]. This is in contrast to the processes that take place on the copper current collector. Aluminium passive films in LIB electrolytes have been extensively studied and characterised via potentiostatic and potentiodynamic electrochemical tests, along with XPS, electrochemical quartz crystal microbalance (EQCM), Fourier-transform infrared spectroscopy, ToF-SIMS and NMR analysis [15, 23, 47, 76 – 80].

In LiPF₆ based electrolytes the passive film is made up of an inner layer of air-formed Al₂O₃ and an outer layer of AlF₃, which significantly contributes to aluminium's corrosion resistance in LIB conditions [19, 23, 27, 75, 76, 81, 82]. The AlF₃ component of the film forms at voltages above a critical potential value, found to be ∼4.1 V and its thickness was observed to increase with applied potential [75]. The critical potential for AlF₃ film formation was also observed to vary from ∼3.8 to 4.2 V depending on the thickness of the pre-existing oxide film [13, 26, 74, 76, 79, 81, 83]. The thickness of the passivation layers has been assessed by electrochemical impedance spectroscopy (EIS) and EQCM—indicating ultra-thin, nanoscale layers are formed: the Al₂O₃ layer is ∼1–2 nm thick and the AlF₃ layer is just ∼1 nm thick [75].

A typical CV of aluminium metal against an Li counter electrode (and Li reference electrode) in an LiPF₆ based electrolyte can be seen in figure 15 [47]. Unlike the CVs seen for copper in the previous section, no cathodic reactions (such as SEI formation) were observed between OCP and 0.3 V. As the potential nears 0 V, a large cathodic current density is observed and is ascribed to the Li–Al alloying process. This process is induced by the deposition of dissolved Al³⁺ (from the air formed surface layer). On the anodic scan, a large current density is observed around 0.6 V. This is ascribed to the corresponding de-alloying process and oxidation of the deposited lithium and aluminium metal to oxide on the surface of the aluminium foil as described in equation (16). The morphological implications (such as surface roughness) of the alloying and de-alloying processes are not discussed. Such surface roughness is of interest to study, as it is likely to make the aluminium surface susceptible to corrosion effects such as crevice corrosion. From ∼1 V until ∼3 V, no significant current density is measured. After ∼3 V two very small anodic peaks are measured at 3.7 V and 4.7 V. These reactions are not reversible and so can be ascribed to the formation of the aluminium passive film.

$$\begin{equation}2{\text{Al}} + 3{\text{L}}{{\text{i}}_2}{\text{O}} \to {\text{A}}{{\text{l}}_2}{{\text{O}}_3} + 6{\text{L}}{{\text{i}}^ + } + 6{{\text{e}}^ - }.\end{equation} \tag{ 16 }$$

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Martha et al carried out potentiostatic tests on aluminium foil in LiPF₆ in EC and DMC, the results of which can be seen in figure 16 [78]. These studies showed that current densities under 3.5 V were around ten times higher than those under 3.0 V, 4.0 V and 5.0 V. This high current density at ∼3.5 V was attributed to the formation of an Al₂O₃ film. Markovsky carried out similar experiments on aluminium at 60 °C [77]. Current oscillations were detected at 4.2 V and 4.8 V, which were ascribed to the formation of passive layers from decomposition of the electrolyte salt. It was shown by XPS that these layers comprised of AlF₃, Al₂O₃ and AlPO₄ species.

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Physical properties of the aluminium were also explored after polarisation in LIB electrolytes. Li et al stored aluminium foils in LiPF₆ in EC and EMC at 4.3 V, 4.6 V, 4.9 V and 5.3 V for one week [23]. Large cracks were observed by transmission electron microscopy (TEM) on the sample that was stored at 5.3 V as can be seen in figure 17. No obvious corrosion or cracks were observed on the samples stored at more negative voltages. It would be of interest to correlate the extent of current collector cracking with long-term LIB electrochemical performance.

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As mentioned previously, a protective layer is formed on aluminium current collectors in LiPF₆ based LIB electrolytes. This layer is made up of an inner layer or air formed Al₂O₃ and an outer passive layer of AlF₃. This layer prevents significant aluminium corrosion from taking place during LIB operation.

The composition and presence of impurities in the electrolyte can influence the formation of the passive layer. For example, Al(OH)₃ was observed via XPS on the surface of an aluminium current collector, which was suggested to have formed due to the presence of water contamination in the electrolyte [80]. The use of LiPF₆ over other electrolyte salts (such as amide salts) has been observed to improve aluminium corrosion resistance during LIB operation [76, 80]. LiPF₆ based electrolytes form fluoride-rich passivating films on the aluminium that increases the long term resistance. Despite this positive effect on aluminium, fluorinated electrolyte salts (especially LiPF₆) decompose to form HF which can contribute significantly to overall LIB degradation (as discussed in section 5). Whilst there is therefore some interest in using electrolytes that do not contain fluorinated salts, it will be important to consider the decreased corrosion resistance of aluminium when optimising such systems [13, 84 – 88].

Electrolyte salt concentration has also been observed to influence the corrosion resistance of the aluminium current collector. Zhang et al performed anodic polarisation tests in solutions of LiPF₆ in EC/DMC, varying the salt concentration from 0.001 to 1 M [75]. It was found that the salt concentration must exceed 0.05 M in order to form a sufficiently protective film on the aluminium.

The presence of HF impurities was found to actually suppress the dissolution of aluminium in some cases [1]. When HF was added to an electrolyte containing LiN(CF₃SO₂)₂, the formation of pits was suppressed on an aluminium electrode. It is thought that the addition of HF facilitates the formation of a more protective AlF₃ passive film. In LiPF₆ based electrolytes, the successful passivation (AlF₃ on the Al₂O₃ surface of the aluminium metal) is found to provide sufficient protection to corrosion attack. Again, the implications of this suppressed corrosion due to the AlF₃ film have not been sufficiently studied, however it would be valuable to understand its contribution to long term LIB electrochemical performance.

The microstructure and surface conditions, including defects, of aluminium influence the corrosion behaviour in LIB systems as has been demonstrated experimentally. For example, it was found by examining aluminium polarisation curves that pre-treating aluminium by heating to 480 °C delayed the onset of corrosion in LiClO₄ in EC/DME [27]. Different aluminium alloys may also corrode differently. Alloy 1145 (99.45% minimum aluminium) behaved similarly to alloy 1100 (99.0% minimum aluminium and 0.12% copper) when cycled 100 times in LiPF₆ in PC/DEC. However, electrochemical impedance spectroscopy indicated that alloy 1145 may have better corrosion resistance upon extended cycling [26]. Iwakura et al performed potentiostatic tests that concluded that the electrochemical stability of aluminium in LiClO₄ in EC/DEC increased with aluminium purity; however, this difference was negligible in LiPF₆ in EC/DEC as the protective effect of the fluorinated passive film dominates [13].

Defects such as mechanical holes and native grooves may form during the aluminium foil manufacturing process. Similar defects may also form during cell assembly. Such defects can act as initiation sites for the pitting corrosion process [89, 90]. Zhang et al used SEM and EDX to propose that pitting on aluminium is also initiated by mechanical damage to the aluminium current collector that may occur during the process of casting the cathode slurry onto the current collector [91]. Yang et al used tweezers to purposely scratch the surface of an aluminium foil [76]. Obvious corrosion was observed with SEM at the scratch site after cyclic voltammetry experiments.

Crevice corrosion has also been observed on aluminium foil when it is coated with a porous cathode [27, 89, 92, 93]. The porosity of the cathode coating enables electrolyte infiltration, ensuring good contact with the electrode's active material. This porosity however, also enables the development of localised occluded sites, and potentially therefore crevice corrosion. Such corrosion can be seen—figure 18 shows an SEM image of an aluminium current collector, coated with an LiFePO₄ cathode and anodically polarised at 100 μA cm⁻² for 36 h in an LiPF₆-based battery electrolyte [92]. Part of the porous cathode has been removed, revealing the corroded surface of the current collector. In this paper, the authors investigated the corresponding degradation of overall battery performance observed upon corrosion of the aluminium foil (as can be seen in figure 19). Power fade and capacity fade have been measured to increase significantly with increasing crevice corrosion for LiNi_0.8Co_0.15Al_0.5O₂ vs. graphite cells cycled at 25 °C [89]. It can be seen that aluminium corrosion correlates directly to the overall degradation in LIBs. Unfortunately such studies that directly investigate the implications of current collector corrosion on LIB performance are limited however it is clear that this understanding will be vital for enabling the development of better batteries.

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