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Lithium Superoxide Hydrolysis and Relevance to Li-O Batteries Hsien-Hau Wang, Yun Jung Lee, Rajeev S. Assary, Chengji Zhang, Xiangyi Luo, Paul C. Redfern, Jun Lu, Young Joo Lee, Do Hyung Kim, Tae-Geun Kang, Ernesto Indacochea, Kah Chun Lau, Khalil Amine, and Larry A Curtiss J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.6b12950 • Publication Date (Web): 17 Apr 2017 Downloaded from http://pubs.acs.org on April 18, 2017

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The Journal of Physical Chemistry

Lithium Superoxide Hydrolysis and Relevance to Li-O2 Batteries Hsien-Hau Wang,1 Yun Jung Lee,2 Rajeev S. Assary,1 Chengji Zhang,1, 4 Xiangyi Luo,3 Paul C. Redfern,1 Jun Lu,3 Young Joo Lee,2 Do Hyung Kim,2 Tae-Geun Kang,2 Ernesto Indacochea,4 Kah Chun Lau,5 Khalil Amine,3 and Larry A. Curtiss1 1

Materials Science Division, Argonne National Laboratory, Argonne Illinois 60439, United States 2 Department of Engineering, Hanyang University, Seoul 133-791, Republic of Korea 3 Chemical Sciences and Engineering Division, Argonne National Laboratory, Argonne Illinois 60439, United States 4 Department of Civil and Materials Engineering, University of Illinois at Chicago, Chicago 60607, United States 5 Department of Physics and Astronomy, California State University, Northridge, California 91330, United States.

Abstract Fundamental understanding of reactions of lithium peroxides and superoxides is essential for the development of Li-O2 batteries. In this context, an investigation is reported of the hydrolysis of lithium superoxide, which has recently been synthesized in a Li-O2 battery. Surprisingly, the hydrolysis of solid LiO2 is significantly different from that of NaO2 and KO2. Unlike KO2 and NaO2, the hydrolysis of LiO2 does not produce H2O2. Similarly, the reactivity of Li2O2 toward water differs from LiO2, in that Li2O2 results in H2O2 as a product. The difference in the LiO2 reactivity with water is due to the more exothermic nature of the formation of LiOH and O2 compared to the corresponding reactions of NaO2 and KO2. We also show that a titration method used in this study, based on reaction of the discharge product with a Ti(IV)OSO4 solution, provides a useful diagnostic technique to provide information on the composition of a discharge product in a Li-O2 battery.

1. Introduction Alkali metal oxides have been the subject of various reviews1-2 and a topic in inorganic chemistry textbooks.

3

The reactivity properties of the alkali metals with oxygen differ with

lithium combustion in air giving Li2O, sodium combustion giving the peroxide Na2O2, and the others (M = K, Rb, Cs) giving superoxides MO2.3 At higher temperatures and pressures Na2O2 1 ACS Paragon Plus Environment


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will take up oxygen to form NaO2. Due to the very electropositive nature of the alkali metals, the known solid alkali metal oxides (M = Li, Na, K, Rb, Cs) hydrolyze readily in water via Reactions 1-33 M2O + H2O 2MOH

(Reaction 1)

M2O2 + 2H2O 2MOH + H2O2

(Reaction 2)

2MO2 + 2H2O O2 + 2MOH + H2O2

(Reaction 3)

Lithium superoxide has only been studied in low temperature (4 K) matrices as a molecule, 4 so its bulk properties have not been studied, including Reaction 3. Recently, we reported on the synthesis of a solid phase of LiO2 in an electrochemical 5

cell. Solid LiO2 has been difficult to synthesize in pure form6 because it is thermodynamically unstable with respect to disproportionation, giving Li2O2 and O2.7,8 We found that crystalline LiO2 could be stabilized in an Li–O2 electrochemical cell by using a graphene oxide based cathode with Ir nanoparticles, which seem to serve as a template for the nucleation and growth of LiO2. Various characterization techniques were used to prove the presence of only LiO2 as discharge product.5 These new results on the synthesis of LiO2 provide an opportunity to study the properties of LiO2 for the first time. In this paper we report on a surprising difference of the reactivity of the solid LiO2 with water compared to that of NaO2 and KO2. This result is due to differences in the hydrolysis reaction pathways for the different alkali metal superoxides. In addition, it is shown how this result can be used as an effective titration test for mixtures of Li2O2 and LiO2 in Li-O2 batteries to determine the fraction of Li2O2 present, which is important for understanding the discharge chemistries in Li-O2 batteries. In Section 2 we describe the experimental method used for assessing the reactivity of LiO2 towards water. In Section 3 we present titration results for LiO2 from Li-O2 electrochemical cells, as well as a possible explanation for why LiO2 shows an unexpected difference in hydrolysis compared to other alkali metal superoxides.

2. Experimental Method Initially a Li-O2 cell with an Ir-rGO cathode identical to that used previously5 to produce a LiO2 discharge product was run for one discharge cycle. The reaction of the LiO2 discharge product with water was investigated using a spectrophotometric method9,10 to assess the amount of hydrogen peroxide produced. In this procedure we first establish a calibration curve by adding 2 ACS Paragon Plus Environment

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incremental amounts of Li2O2 to the test reagent Ti(IV)OSO4 (Aldrich, ~2 wt% in sulfuric acid) as shown in Figures 1A and 1B. The discharge product LiO2 on the Ir-rGO cathode is then reacted with the acidic solution of Ti(IV)OSO4 after removal of the electrolyte, but with the cathode material still present. Under the experimental conditions, Ti(IV) exists as TiO2+(aq) complexed with SO42- ion in 1.0 M H2SO4. This method has been used to distinguish hydrogen peroxide from organic peroxides while avoiding interference from dissolved oxygen at low peroxide levels.9 If hydrogen peroxide is present, a color change to yellow/orange occurs due to the formation of a titanium peroxide complex, TiO2SO4. This technique has been used previously for detecting formation of Li2O2,10-11 NaO2,12 and KO213 in M-O2 electrochemical cells. Since the technique uses a strongly acidic solution, any LiOH resulting from reaction 3 is neutralized by the protons. The acidity of the solution was found not to significantly affect the H2O2 detection.9

3. Results and Discussion An acidic solution of Ti(IV)OSO4 was added to the LiO2 discharge product formed in the electrochemical cell to detect hydrogen peroxide. The titration was done using the experimental procedure described in Section 2. When the acidic solution is added, the liquid remained colorless and no notable UV-Vis intensity change was observed as shown in Figure 1C. From the UV-Vis spectra in Figure 1D, the absorbance intensity of 0.04 at 410 nm is essentially negligible based on our calibration curve, which indicates the absence of hydrogen peroxide. In addition, gas bubbles were observed to be released during the procedure. The LiO2 results are in contrast to the case for KO2 and NaO2 where an obvious color change is observed for this procedure, consistent with the presence of a hydrogen peroxide product from Reaction 3.12,13 It is also in contrast to the case of Li2O2 where the same titration test is also positive for the presence of H2O2 (Reaction 2) and a color change is observed.14 The possibility that the cathode support (Ir-rGO) for the LiO2 formed in the electrochemical cell could decompose the hydrogen peroxide was also investigated and could not account for the absence of H2O2 (see supplemental information). The free energies for hydrolysis of MO2 to MOH, H2O2, and O2 [Reaction (3)] for M = Li, Na, K were computed from literature values to gain insight into the differing experimental results for these three alkali metals. The reaction energies are given in Table 1 and are for MO2, M2O2, and MOH in their solid state and the other species in liquid (H2O, H2O2) and gaseous (O2) 3 ACS Paragon Plus Environment


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states. Also given in Table 1 is the energy for Reaction 1 (hydrolysis of M2O) and Reaction 2 (hydrolysis of M2O2), which are based on the solid states of M2O, M2O2, and MOH. The reaction free energies in Table 1 are approximate and meant only to give qualitative trends for the different alkali metals as the energies neglect factors such as crystallization, dissolution, and solvation energies involved in these solution phase reactions. The effect of the neglect of these additional factors is probably the reason why most of the hydrolysis reactions in Table 1 are endergonic despite the fact that the superoxides and peroxides are known to readily hydrolyze.3 Hydrolysis of the M2O2 peroxides (Table 1) to the corresponding MOH hydroxides plus H2O2, [Reaction 2], is endergonic for lithium (+55 kJ/mol), sodium (+45 kJ/mol), and potassium (+22 kJ/mol), based only on solid state free energies. Likewise, Reaction 3 is endergonic for sodium (32 kJ/mol) and potassium (75 kJ/mol) superoxides, and exergonic for lithium superoxide (-20 kJ/mol). The effect of inclusion of water molecules on these reactions based on molecular states was calculated using high level quantum chemical15,16 calculations. The results show a significant shift from being unfavorable in the gas phase to being favorable when solvation effects are included (see Table S1). This effect is consistent with both the observed hydrolysis of these solids3 and the use of these reactions for quantification of Li2O2, NaO2, and KO2 by titration.12,13,14 A comprehensive study of solvation effects for the hydrolysis reactions of peroxides and superoxides is beyond the scope of this study, but the qualitative trends from the energies in Table 1 are useful for explaining the results in this study. In order to understand the titration results in this paper we also need to consider the reaction for complete hydrolysis of MO2. In this case the reaction proceeds completely to MOH without the H2O2 formation 4MO2 + 2H2O 3O2 + 4MOH

(Reaction 4)

Thus, the hydrolysis of MO2 superoxides to MOH can occur via partial hydrolysis (Reaction 3) or via complete hydrolysis (Reaction 4). A complete hydrolysis of MO2 to MOH (Reaction 4) could occur by a step-wise mechanism consisting of (1) Reaction 3 described above and (2) subsequent reaction with H2O2 4MO2 + 2H2O 2MOH + H2O2 + O2 + 2MO2 3O2 + 4MOH.

(Reaction 5)

A schematic of the potential energy surface for a two-step mechanism of complete hydrolysis of superoxide is shown in Figure 2 4 ACS Paragon Plus Environment

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The qualitative trends in the reaction energies in Table 1, including that for Reaction 4 above, can be used to help understand why the hydrolysis of Li is different from that of Na and K. Based on the much more favorable energies for Reactions 3 and 4 for LiO2 compared to NaO2 and KO2, it is likely that LiO2 follows a concerted reaction as shown in Figure 2, to form LiOH and O2. A concerted reaction is a chemical reaction in which all bond breaking and bond making occurs in a single step and reaction intermediates are not involved. In a concerted reaction the H2O2 intermediate shown in Reaction 5 will not be present in the mechanism for LiO2 because of the large exergoncity of the reaction (Table 1). The gas evolution observed during the titration experiment for LiO2 is consistent with Reaction 4. In contrast NaOH and KOH formation from NaO2 and KO2, respectively, likely proceed via a stepwise mechanism instead of a concerted mechanism because their reactions are not as exergonic. Thus, this allows the H2O2 product in Reaction 3 to exist long enough to be quantified.12,13. A stepwise mechanism with a kinetically controlled second step for NaO2 and KO2 is consistent with the need to use a catalyst for complete hydrolysis of NaO2 and KO2 (Reaction 4).17 In order to further confirm that the titration technique used is correct for the LiO2 discharge product, we have tested for the expected LiOH product based on Reaction 3 or 4. This was done by pH measurement of a solution after addition of known quantity of H2O is added to a cathode after one discharge cycle in a Li-O2 cell with an Ir-rGO cathode under the same conditions as described for the titration. The resulting solution was strongly basic with a pH of 11.84, consistent with the presence of OH- anion from LiOH. In addition, analysis of the current used and the discharge time indicated that the amount of OH- produced is 95% of that expected based on Reaction 3 or 4 (see supplement information for more information). Thus, this result and the lack of the presence of H2O2 from the Ti(IV)OSO4 titration indicates that hydrolysis of LiO2 occurs by Reaction 4. These new results on the hydrolysis of LiO2 provide a useful method to determine the amount of Li2O2 present in the discharge product, which may be composed of a mixture of components including LiO2. In addition, the pH analysis can be used to confirm the presence of LiO2 and quantitatively measure the LiOH produced, which correlates with the amount of LiO2. Table 2 compares titration results based on the Ti(IV)OSO4 technique for three different Li-O2 cathode materials including (1) the Ir-rGO cathode that has been confirmed to give only LiO2 by several other techniques,

5

(2) a carbon paper cathode that has been confirmed to give only 5 ACS Paragon Plus Environment


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Li2O2,14 and (3) an activated carbon cathode that gives a mixture of LiO2 and Li2O2.

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17-19

The

titration results in Table 2 are consistent with the characterization of the discharge product for all three cathode materials. The Ir-rGO titration result as mentioned earlier gave a negative result consistent with the presence of LiO2. The work of Wu et. al. 14 for the carbon paper cathode gave a positive result for 100% Li2O2. The titration for the activated carbon discharge product indicated about 50% of the discharge product was Li2O2. This is consistent with a kinetics study on the same system that gave about 50% Li2O2.

17

We also note that the uncertainty in the

titration results is about +/-3% based on the potential error in the measurement of the mass for the calibration curve (Figure 1).

4. Conclusions The following conclusions can be drawn from this study of the hydrolysis of lithium superoxide, which has recently been synthesized in a Li-O2 battery: (1) The hydrolysis of solid LiO2 is significantly different than that of NaO2 and KO2. Unlike KO2 and NaO2, the hydrolysis of LiO2 does not give hydrogen peroxide as a product. Similarly the reactivity of Li2O2 differs from LiO2, in that the former results in H2O2 as a product. (2) The reason for the difference in the LiO2 reactivity with water is likely because lithium superoxide follows a concerted reaction to form LiOH and O2 due to the more strongly exergonic nature of the reaction compared to for NaO2 and KO2. In a concerted reaction the H2O2 intermediate will not occur as it does for NaO2 and KO2, or for Li2O2. (3) The titration method, based on the Ti(IV)OSO4 technique, applied here provides a useful method to determine the amount of Li2O2 present in a discharge product that may be composed of a mixture of other components such as LiO2. In addition, a pH analysis can be used to confirm the presence of LiO2 and quantitatively measure the LiOH produced. These new results on the hydrolysis properties of LiO2 will be of use in understanding and characterizing reactions that occur in Li-O2 batteries including the presence of LiO2 in the discharge product and the effect of presence of water during reactions. Acknowledgement

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This work was supported by the U.S. Department of Energy under Contract DE-AC0206CH11357 from the Vehicle Technologies Office, Department of Energy, Office of Energy Efficiency and Renewable Energy. We acknowledge grants of computer time on the ALCF Fusion and Blues Clusters at Argonne National Laboratory. Supporting Information Available: Calibration data for titration measurement; molecular reaction energies; analysis of LiOH product during titration. References 1. Cotton, F. A., Progress in Inorganic Chemistry; John Wiley and Sons: New York, 1962; Vol. 4, p 125-197. 2. Borgstedt, H. U.; Mathews, C. K., Applied Chemistry of the Alkali Metals; Plenum Press: New York, 1987. 3. Cotton, F. A.; Wilkinson, G., Advvanced Inorganic Chemistry; Wiley-Interscience: New York, 1988. 4. Andrews, L., Matrix Infrared Spectrum and Bonding in the Lithium Superoxide Molecule, LiO2. J. Amer. Chem. Soc. 1968, 90, 7368-7370. 5. Lu, J.; Lee, Y. J.; Luo, X.; Lau, K. C.; Asadi, M.; Wang, H.-H.; Brombosz, S.; Wen, J.; Zhai, D.; Chen, Z.; Miller, D. J.; Jeong, Y. S.; Park, J.-B.; Fang, Z. Z.; Kumar, B.; Salehi-Khojin, A.; Sun, Y.-K.; Curtiss, L. A.; Amine, K., A Lithium–Oxygen Battery Based on Lithium Superoxide. Nature 2016, 529, 377–382. 6. Sangster, J.; Pelton, A. D., The Li-O (Lithium-Oxygen) System. Journal of Phase Equilibria 1992, 13, 296–299. 7. Lau, K. C.; Curtiss, L. A.; Greeley, J., Density Functional Investigation of the Thermodynamic Stability of Lithium Oxide Bulk Crystalline Structures as a Function of Oxygen Pressure. J. Phys. Chem. C 2011, 115, 23625-23633. 8. Kang, S.; Mo, Y.; Ong, S. P.; Ceder, G., A Facile Mechanism for Recharging Li2O2 in Li−O2 Batteries. Chem. Mater. 2013, 25, 3328-3336. 9. Sellers, R. M., Spectrophotometric Determination of Hydrogen Peroxide Using Potassium Titanium( Iv) Oxalate. Analyst 1980, 105, 950-954. 10. O'Sullivan, D. W.; Tyree, M., The Kinetics of Complex Formation between Ti(Iv) and Hydrogen Peroxide. Int. J. Chem. Kinet. 2007, 39, 457-461. 11. Eisenberg, G. M., Colorimetric Determination of Hydrogen Peroxide. Ind. Eng. Chem., Anal. Ed. 1943, 15, 327-328. 12. Hartmann, P.; Bender, C. L.; Sann, J.; Du¨rr, A. K.; Jansen, M.; Janek, J. r.; Adelhelm, P., A Comprehensive Study on the Cell Chemistry of the Sodium Superoxide (NaO2) Battery. Phys. Chem. Chem. Phys. 2013, 15, 11661. 13. Ren, X.; Lau, K. C.; Yu, M.; Bi, X.; Kreidler, E.; Curtiss, L. A.; Wu, Y., Understanding Side Reactions in K−O2 Batteries for Improved Cycle Life. ACS Appl. Mater. Interfaces 2014, 6, 19299-19307. 14. Yu, M.; Ren, X.; Ma, L.; Wu, Y., Integrating a Redox-Coupled Dye-Sensitized Photoelectrode into a Lithium–Oxygen Battery for Photoassisted Charging. Nat. Commun. 2014, 5, 5111. 15. Curtiss, L. A.; Redfern, P.; Raghavachari, K., Gaussian-4 Theory. J. Chem. Phys. 2007, 126, 084108. 16. Curtiss, L. A.; Redfern, P.; Raghavachari, K., Gaussian-4 Theory Using Reduced Perturbation Orders. J. Chem. Phys. 2007, 127, 124105. 7 ACS Paragon Plus Environment


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17. Zhai, D.; Wang, H.-H.; Yang, J.; Lau, K. C.; Li, K.; Amine, K.; Curtiss, L. A., Disproportionation in Li−O2 BaƩeries Based on a Large Surface Area Carbon Cathode. J. Amer. Chem. Soc. 2013, 135, 15364 15372. 18. Zhai, D.; Lau, K. C.; Wang, H.-H.; Wen, J.; Miller, D. J.; Lu, J.; Kang, F.; Li, B.; Yang, W.; Gao, J.; Indacochea, E.; Curtiss, L. A.; Amine, K., Interfacial Effects on Lithium Superoxide Disproportionation in Li-O2 Batteries. Nano Lett. 2015, 15, 1041-1046. 19. Zhai, D.; Wang, H.-H.; Lau, K. C.; Gao, J.; Redfern, P. C.; Kang, F.; Li, B.; Indacochea, E.; Das, U.; Sun, H.-H.; Sun, H.-J.; Amine, K.; Curtiss, L. A., Raman Evidence for Late Stage Disproportionation in a Li−O2 BaƩery. J. Phys. Chem. Lett. 2014, 5, 2705−2710.


A

2 Abs Abs Abs

1.5

Absorbance

Abs Abs Abs Abs

1

Blank 0.1 mg 0.7 mg 1.3 mg 1.8 mg 2.4 mg 3.2 mg

2

B

y = 0.17977 + 0.47708x R= 0.99652

1.5

0.5

1

0.5

0

0

400

500

600

700

800

0

0.5

1

C

1.5

2

2.5

3

3.5

Li O Wt mg

Wavelength (nm)

2

D

2

0.2 Abs Abs Li_discharge Abs_Li_dis-sept

0.15

Absorbance

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Absorbance

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Blank Discharge cathode Discharge cathode + separator

0.1

0.04950

0.05

0.03667

0

0.00633

400

500

600

700

800

Wavelength (nm) Figure. 1 (A) UV-Vis measurements (dual beam) of Li2O2 powders (Aldrich) added to 50 mL of TiOSO4 reagent (aq); (B) Calibration curve for Li2O2 in TiOSO4 ; (C) Photograph of the cathode (Ir-rGO/LiO2) discharged to 1000 mAh/g, dried under Ar for 1 hr, and then soaked in 3.00 mL of TiOSO4(aq); no apparent color change occurred upon titration indicating no Li2O2 is present and also gas bubbles were observed coming off the liquid; (D) Plot of the absorbance of the liquid from the discharged cathode in C; the absorbance is near negligible from the calibration.



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Figure 2 Schematic of hypothetical potential energy pathways for stepwise and concerted MO2 reactions. The A B reaction corresponds to Reaction 3 and AC corresponds to Reaction 4 in the text. The energy schemes are approximate and assume solution phase effects not included in Table 1.


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Table 1: Estimated free energies (kJ/mol) at 298 K of Reactions 1-4 from experimental data.a ∆G, kJ/mol Reaction M= Li M=Na M=K M2O + H2O 2MOH Reaction 1 -79.6 -146.1 -198.1 M2O2 + 2H2O 2MOH + H2O2 Reaction 2 54.9 44.8 21.7 2MO2 + 2H2O 2 MOH + H2O2 + O2 Reaction 3 -20.0 32.0 75.4 4MO + 2 H O 4MOH + 3 O Reaction 4 -274.3 -169.6 -82.8 2 2 2 a All reaction energies are derived from data in Lange's Handbook Of Chemistry 15th edition McGraw Hill: New York, 1999, except for LiO2, which is a theoretical value from Ref. 7. The

theoretical value is used for the free energy of LiO2 because there is no experimental value for this quantity The theoretical value is from a first principles density calculations and should be quite accurate. The ∆G’s do not include solution phase reaction effects. Table 2: Summary of titration results for three Li-O2 discharge products with different compositions. Cathode Material Ir-rGO (D) Activated carbon (D) Carbon paper P50

Color change No Yes

Amount of Li2O2 0% ~50%

Yes

100%

Gas evolution Yes Not noticeable ----

Other Evidence for product composition DEMS: 1e/O2 (Ref. 5) Kinetics data: ~50% Li2O2 (Ref. 18) Mass/electron balance: 100% Li2O2 (Ref. 14)

Reference This work This work Ref. 14



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Table of Contents Figure


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Table of Contents Figure

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