AlâAir Batteries - American Chemical Society

Letter pubs.acs.org/JPCL

Al−Air Batteries: Fundamental Thermodynamic Limitations from First-Principles Theory Leanne D. Chen,†,‡ Jens K. Nørskov,†,‡ and Alan C. Luntz*,† †

SUNCAT Center for Interface Science and Catalysis, SLAC National Accelerator Laboratory, 2575 Sand Hill Road, Menlo Park, California 94025, United States ‡ Department of Chemical Engineering, Stanford University, Stanford, California 94305-5025, United States S Supporting Information *

ABSTRACT: The Al−air battery possesses high theoretical specific energy (4140 W h/ kg) and is therefore an attractive candidate for vehicle propulsion. However, the experimentally observed open-circuit potential is much lower than what bulk thermodynamics predicts, and this potential loss is typically attributed to corrosion. Similarly, large Tafel slopes associated with the battery are assumed to be due to film formation. We present a detailed thermodynamic study of the Al−air battery using density functional theory. The results suggest that the maximum open-circuit potential of the Al anode is only −1.87 V versus the standard hydrogen electrode at pH 14.6 instead of the traditionally assumed −2.34 V and that large Tafel slopes are inherent in the electrochemistry. These deviations from the bulk thermodynamics are intrinsic to the electrochemical surface processes that define Al anodic dissolution. This has contributions from both asymmetry in multielectron transfers and, more importantly, a large chemical stabilization inherent to the formation of bulk Al(OH)3 from surface intermediates. These are fundamental limitations that cannot be improved even if corrosion and film effects are completely suppressed.

R

not react violently with water), which could also provide all of the advantages without the safety concerns. Al, being the most abundant metal in earth’s crust and possessing very low molecular mass, is an attractive candidate for these aqueous metal−air batteries. The Al−air system is described by the overall reaction Al + 3/4O2 + 3/2H2O → Al(OH)3 (U0 = 2.70 V), and its MTSE is 4140 W h/kg including the mass of water.3 Thus, the Al−air battery has considerable potential for improvement in practical specific energy and hence EV range over that possible with a Li-ion battery. Moreover, the Al−air battery is capable of outputting high current densities and therefore could be used in highpower applications.4 Concentrated alkaline (NaOH/KOH) electrolytes are most commonly used with this system.4,5 The discharge product, Al(OH)3, is an intermediate in the Bayer process, the principal industrial process of Al metal production, and therefore, an infrastructure already exists for recycling the battery materials.6 Furthermore, although the aqueous Al−air battery is not electrically rechargeable, given the existing extensive Al refining/recycling economy, a primary battery would be convenient for the consumer as entire battery packs could be replaced quickly at refilling stations. We first summarize the current understanding of aqueous bulk electrochemistry relevant to the Al−air battery with the

ising worldwide energy demand, coupled with climate change associated with usage of fossil fuels, motivates the search for alternative, renewable energy sources. In particular, the widespread implementation of an electrified transportation system is highly desirable because the majority of petroleum, which accounts for one-third of the worldwide primary energy source, is used in vehicle propulsion. Current state-of-the-art electric vehicles (EVs) principally utilize the Li ion battery, which is the battery possessing the highest practical specific energy that has been sufficiently developed to be feasible for transportation. However, the aforementioned practical specific energy is still very low (∼200 W h/kg) compared to that of gasoline in an internal combustion engine (∼1300 W h/kg), and thus, current EVs are greatly limited by their range. In addition, large format Li-ion batteries are presently prohibitively expensive for mass-market adoption of EVs and have safety concerns related to using flammable nonaqueous electrolytes. In principle, metal−air batteries could circumvent these limitations as they possess much larger maximum theoretical specific energies (MTSEs) due to eliminating the necessity of cathode material storage, and they are projected to be less costly than Li-ion batteries. The nonaqueous Li−air battery has received a great deal of attention over the past few years due to its extremely high MTSE, 11680 W h/kg. However, it currently has many technical limitations1,2 and similar safety concerns to those of the Li-ion battery due to the use of flammable electrolytes. An alternative is to use aqueous metal−air batteries with compatible metals (i.e., metals that do © XXXX American Chemical Society

Received: November 15, 2014 Accepted: December 17, 2014

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dx.doi.org/10.1021/jz502422v | J. Phys. Chem. Lett. 2015, 6, 175−179

The Journal of Physical Chemistry Letters

Letter

following half-reactions, where all potentials are given versus the standard hydrogen electrode (SHE). In acid, formation of the aqueous metal ion dominates Al(s) ⇌ Al3 +(aq) + 3e−

U0 = −1.67 V

(1)

In basic media, the favored processes are formation of the bulk hydroxide Al(s) + 3OH−(aq) ⇌ Al(OH)3(s) + 3e−

U0 = −2.30 V (2)

and the aluminate anion Al(s) + 4OH−(aq) ⇌ Al(OH)4 −(aq) + 3e− U0 = −2.31 V

(3)

Parasitic water reduction, where different areas of the Al electrode act as both cathodic (H2 evolution) and anodic sites (Al oxidation) creating an internal short-circuit on the anode, may also occur 1 H 2O(l) + e− ⇌ OH−(aq) + H 2(g) U0 = −0.83 V 2

Figure 1. Bulk Pourbaix diagram outlining the anode and cathode reactions in the Al−air battery. All reactions except that marked cathode occur at the Al anode.

surface.12,13 However, it must be emphasized that the standard potentials in eqs 1−5 are referenced from thermodynamic bulk formation energies and thus only describe the initial and final states in redox processes without any bearing on the energetics of the surface mechanism by which they actually occur in a battery. Therefore, we must obtain the energies of all surface intermediates in the electrochemical dissolution of Al metal in alkaline conditions in order to gain insight into the potential that can actually be yielded by the Al surface in these Al−air batteries. Here, we present, to the best of our knowledge, the first ab initio study on the stepwise hydroxide-assisted mechanism for anodic Al dissolution in alkaline media involving single-electron transfers and subsequent formation of bulk Al(OH)3. This mechanism has been proposed previously by several experimental groups9,12,13 to describe the electrochemistry of Al−air batteries; moreover, a similar theoretical methodology has been applied in many studies of electrocatalysis7,14 as well as recent studies on the aqueous Zn−air battery15 and the nonaqueous Li−air battery.16 We use density functional theory (DFT) to calculate the free energies of all surface intermediates; see the Supporting Information (SI) for details of the ab initio calculations. Note that in the following, we neglect any potential drops associated with the air cathode by focusing entirely on the Al anode relative to SHE. We emphasize that in this study, we are interested in answering a fundamental question: what is the limiting potential that can be extracted from the Al anode surface in the absence of parasitic reactions? We will demonstrate that both open-circuit and moderate current density potentials calculated by DFT are similar to experimental values in the absence of corrosion and film formation. Our results suggest that the difference between the potential predicted by bulk thermodynamics and one obtained in the stepwise mechanism arises from (1) asymmetry in free energies of intermediates in multielectron transfer reactions and, more importantly, (2) the additional chemical stabilization upon formation of bulk Al(OH)3 from surface-adsorbed Al(OH)3 units. These data imply that the OCP of this electrode is only around −1.87 V versus SHE at pH 14.6, rather than the widely cited −2.34 V, and according to the current model, this OCP cannot be improved by eliminating parasitic

(4)

Note that this corrosion reaction also consumes Al but occurs separately from the productive dissolution process that provides electrons to the external circuit. Finally, the cathodic oxygen reduction reaction in base during discharge is given by 1 O2(g) + H 2O(l) + 2e− ⇌ 2OH−(aq) U0 = 0.40 V 2 (5)

The air cathode has a minimum overpotential of ∼0.4 V7 and therefore contributes ∼0 V to the overall battery potential at pH 14. This implies that, conveniently, all anode potentials referenced to SHE are approximately the negative of the full Al−air battery potentials. The potentials from eqs 1−5 are referenced from thermodynamic reaction energies, where all compounds are in their standard states.3 For solvated species, this is defined to be at a concentration of 1 M, and this implies that all potentials involving OH− in the reaction are given at pH 14. The above electrochemical equations are plotted in the bulk Pourbaix diagram in Figure 1 for ease of visualization using the Nernst equation to account for pH effects on the potential for all reactions with n electrons involving OH− 0.0592 U = U0 + (14 − pH) (6) n Despite the aforementioned favorable properties, the experimentally determined open-circuit potential (OCP) of high-purity Al (99.999%) has been reported to be only ∼−1.6 V8,9 versus SHE in 4 M hydroxide electrolyte (pH = 14.6), compared to a theoretical OCP of −2.34 V versus SHE at this pH. The mixed potential resulting from parasitic water reduction in eq 4 is believed to be the dominant contributor to the difference between the thermodynamic and observed potentials in alkaline conditions.4,10,11 In fact, using Al alloys or aqueous inhibitors that quench much of the Al corrosion gives a more negative OCP of ∼−1.8 V versus SHE.4,8 In addition, at higher current densities (hundreds of mA/cm2), the output potential becomes significantly less negative due to a surprisingly large Tafel slope of 300−500 mV/decade.9,12,13 This large Tafel slope has traditionally been described as arising from film (Al(OH)3 and Al2O3) formation on the Al 176



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reactions. We also show that high Tafel slopes are anticipated by the surface mechanism. We outline the following stepwise mechanism for hydroxideassisted aluminum oxidation because single electron transfers are much more probable than coherent multielectron transfers and adsorptions. Moreover, in order to dissolve an Al atom from an electrode surface, the bonds to its neighbors must be weakened, which would be energetically unfavorable without stabilization from another source, that is, solvent/solute molecules forming bonds to an Al atom.17 Al + OH−(aq) ⇌ AlOH* + e−

(7)

AlOH* + OH−(aq) ⇌ Al(OH)2 * + e−

(8)

Al(OH)2 * + OH−(aq) ⇌ Al(OH)3 * + e−

(9)

Al(OH)3 * + OH−(aq) ⇌ Al(OH)4 −(aq)

(dissolution)

Figure 2. Free-energy diagram of stepwise Al anodic dissolution on the stepped surface at 0 V versus SHE (black) and the limiting potential for this facet, −1.87 V versus SHE (blue). The dashed line represents the calculated free energy of formation for bulk Al(OH)3 as a reference.

(10)

Al(OH)4−(aq)

⇌ Al(OH)3(s) +

OH−(aq)

(precipitation) (11)

Here, Al denotes a surface metal atom, (aq) refers to a solvated species, (s) refers to a bulk solid, and * denotes a surfaceadsorbed species. Equations 7−9 are electrochemical processes and describe single-electron transfers from the metal to external circuit concerted with OH− adsorption (see the SI for details on the treatment of solvated OH− in our DFT calculations). Equation 10 represents detachment of a single molecular Al(OH)3 unit from the electrode surface and forming the aluminate anion in solution, and eq 11 describes precipitation of solvated Al(OH)4− to yield the final product, bulk Al(OH)3, spatially apart from the electrode surface. Because the chemical potential of the aluminate anion depends on its concentration in the solvent whereas the bulk Al(OH)3 chemical potential stays constant, the two species will eventually reach equilibrium, and we show only bulk Al(OH)3 in the freeenergy diagrams for simplicity. Pertaining to the accuracy of these DFT calculations, we note that the calculated free energy of formation for bulk Al(OH)3 at pH 14.6 is −6.94 eV, close to the tabulated thermodynamic value of −7.02 eV.3 The calculated lattice constant of bulk Al (4.05 Å) is also very close to the experimentally determined lattice constant (4.04 Å).18 Furthermore, the average error in chemisorption energy associated with the RPBE exchange−correlation functional used in this study is 0.2 eV.19 Finally, we only consider dissolution on surfaces free of adsorbate coverages because (1) we find that multiple OH* adsorption to the same Al atom is more stable than adsorption on different Al atoms (see the SI for relevant results) and (2) once an adsorbed Al(OH)3* forms, dissolution is thermodynamically favorable; therefore, an OH-free surface (with the exception of the active site) likely dominates throughout the dissolution process in the current thermodynamic picture. Note that this is the thermodynamic implication from our initial postulate not to consider Al dissolution from film-covered surfaces. On the basis of eqs 7−11, we calculate free energies of intermediates AlOH*, Al(OH)2*, and Al(OH)3* on three different Al facets, close-packed/111 (see Figure 3), stepped/ 211 (see Figure 2), and kinked (see Figure S3 in the SI). The ΔG of each electrochemical step is a function of the potential (see the SI for a detailed discussion of how potential effects are incorporated in the DFT results); as the potential is shifted

Figure 3. Free-energy diagram of stepwise Al anodic dissolution on the close-packed surface at 0 V versus SHE (black) and the limiting potential for this facet, −1.47 V versus SHE (blue). The dashed line represents the calculated free energy of formation for bulk Al(OH)3 as a reference.

toward more negative values, the energy of the electrons is raised according to ΔG = −neU, and the ΔG of each electrochemical step is shifted in the positive direction (less favorable). Thus, there exists a lowest potential at which all steps are either thermoneutral or exergonic. This is called the limiting potential because if we attempt to shift the potential toward even more negative values, at least one step in the mechanism would be endergonic, that is, have a thermodynamic barrier. Only sites with the largest adsorption energies are active (undercoordinated atoms at steps and kinks) at open circuit, whereas adsorption on the terrace will also contribute significantly to the discharge potential at higher current densities because the latter represents the dominant structure of the surface.20 Therefore, we describe the OCP with adsorption at stepped (Figure 2) and kinked (see the SI) sites. The calculated limiting potential at a step is −1.87 V, the computed maximum OCP that can be harnessed from the Al anode, and is similar to the experimentally determined OCP 177



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crystallization, and this energy cannot be harnessed electrochemically. The Al electrode, in the absence of corrosion, exhibits a large deviation from the commonly assumed thermodynamic potential to form bulk Al(OH)3. This is in part due to an electronic factor, that is, the asymmetry between the individual electrochemical steps as the bonds to the surface increase in strength with the number of adsorbates (∼0.2 eV per adsorbate, or 3% of the total energy). More importantly, the majority (14−29% of the thermodynamic formation energy depending on the Al facet) of this deviation results from the numerous additional interactions of Al(OH)3 units in the bulk, which are chemical in nature. This is an intrinsic property that arises from the stepwise nature of the surface mechanism and must be clearly distinguished from the overall bulk process. Therefore, according to the present DFT results, the maximum OCP of the Al electrode is only around −1.87 V versus SHE at pH 14.6 rather than the widely cited −2.34 V. The implications of this conclusion are profound for Al−air batteries; while inhibiting corrosion may affect the coulombic efficiency of the battery, such efforts can do little to increase the output potential of the battery above 1.9 V. Note that when significant Al corrosion is present, even lower output potentials are obtained experimentally.4 In addition, comparing the surface thermodynamics at the step/kink to terrace sites also implies a high Tafel slope, even in the absence of film formation, and this ultimately limits the maximum power obtainable from the Al−air battery.

values in the absence of corrosion. An unfavorable step vacancy formation energy is included in the last step of the dissolution scheme. When a higher current needs to be drawn from the battery, terrace atoms necessarily start to dissolve because they make up a larger proportion of the overall surface. Therefore, we model the battery under discharging conditions with the close-packed surface shown in Figure 3. Again, the last step in the dissolution mechanism includes an unfavorable vacancy formation energy. The calculated limiting potential is −1.47 V, ∼0.4 V greater than at the step and kink. As more coordinated surface Al atoms bind less strongly to adsorbates, the limiting potential is less negative than that on the stepped surface. The large shift in the potential required for the minimal (step and kink) and larger (terrace) current densities implies a very large Tafel slope on the order of several hundred mV/decade.20 In the following, we will address why the Al anode yields substantially lower potentials than what bulk thermodynamics predicts. In reactions involving n electron transfers, the observed potential is equal to its theoretical value if and only if both of the following conditions hold: (1) the entirety of the energy released is electrochemical in nature and (2) all electrochemical steps release exactly 1/n of the total energy. As Figures 2 and 3 illustrate, neither condition is fulfilled in the case of the Al anode. The second and third adsorptions of OH− each gain more energy (∼0.2 eV, or 3% of the total thermodynamic energy) than the previous step. Moreover, the last step in the mechanismdissolution of Al(OH)3* from the surface to form Al(OH)3 (s)corresponds to an energy gain of 0.94 and 2.0 eV (14 and 29% of the total thermodynamic energy), respectively, for the stepped and close-packed surfaces. As the last step is completely chemical in nature, this energy is not available in the electrochemical process and is released as wasteful thermal energy. Therefore, the maximum OCP that can possibly be yielded by the Al anode is only −1.87 V versus SHE according to the present DFT results. We would like to further address the origin of the large energy difference between adsorbed Al(OH)3* and bulk Al(OH)3 by examining the structures of both species. In Figure 4, we see that each Al center interacts with three OH groups in a trigonal geometry for surface-adsorbed

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ASSOCIATED CONTENT

S Supporting Information *

Details of the electronic structure calculations and derivation of the computational hydrogen electrode (CHE) that incorporates potential effects and treatment of solvated OH− are available. Geometries of the most stable pure water bilayer and hydroxide/water bilayer on Al(111), free energy diagram of discharge on the kinked surface, and adsorption energies on different sites are also included. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors gratefully acknowledge support from the Office of Basic Energy Sciences of the U.S. Department of Energy to the SUNCAT Center for Interface Science and Catalysis. L.D.C. acknowledges financial support from the Natural Sciences and Engineering Research Council of Canada for the CGS-D3 fellowship and the ReLiable project (#11-116792) funded by the Danish Council for Strategic Research. We also thank Dr. S. Siahrostami for helpful discussions.

Figure 4. (Left) A single surface-adsorbed Al(OH)3* unit. (Right) Local environment of one Al(OH)3 unit in the bulk.

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Al(OH)3*. However, in the bulk, each Al center actually interacts with six OH groups in an octahedral geometry, as well as three other Al atoms in a trigonal geometry. The calculated metal−metal distance is 2.961 Å, compared to 2.865 Å in the bulk metal. In the dissolution mechanism, the electrochemical processes yield the structure on the left, which does not describe the local environment of Al(OH)3 in the bulk. The additional interactions stabilize each Al(OH)3 unit upon

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AlâAir Batteries - American Chemical Society

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