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B(SCN) : A New Weakly Coordinating Anion in the Tetracyanoborate Family Mingmin Zhong, Hong Fang, and Puru Jena J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.7b10332 • Publication Date (Web): 11 Dec 2017 Downloaded from http://pubs.acs.org on December 17, 2017

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The Journal of Physical Chemistry

B(SCN)4–: A New Weakly Coordinating Anion in the Tetracyanoborate Family Mingmin Zhong1, †, Hong Fang2, †, *, Puru Jena2, * 1

School of Physical Science and Technology, Southwest University, Chongqing 400715, China.; 2

Department of Physics Virginia Commonwealth University Richmond, Virginia 23284, USA.

† These are co-first authors. *[email protected] (H. F.) *[email protected] (P. J.)

Abstract Halogen-free

electrolytes

are

important

for

the

development

of

the

next-generation rechargeable Li ion batteries. Due to its ultra-high oxidative stability and weak binding with Li+, the bison anion B(CN)4– has recently been found to be a potential candidate for such electrolytes. Unfortunately, LiBison salt has very poor solubility in ionic liquids, leading to the possible formation of large aggregates, thus hindering its potential for commercial application. Here, by replacing the cyanide with much less toxic thiocyanide, we propose a new weakly-coordinating anion, B(SCN)4–, that does not suffer from this drawback. We show that this new member of the tetracyanoborate family is unlikely to form aggregated complexes with Li+ and would be a good candidate to be used in electrolytes.

1

ACS Paragon Plus Environment

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INTRODUCTION The development of rechargeable batteries based on metal ions is vital to our energy future.1-6 As one of the key components of a battery, an electrolyte produces ionically conductive solution when dissolved in a solvent. A good electrolyte should show high solubility in solvents and produce less reactive and less toxic solutions. Recently, halogen-free cluster ions have found their applications in energy materials.7-16

These

cluster

ions

are

called

hyper/super-halogens

or

pseudohalogens.17-19 The former are characterized by vertical detachment energies (VDE) that are higher than those of elementary halogens, and the latter have VDEs comparable to those of halogens. Many halogen-free electrolytes showing favorable properties can be generated from these cluster ions.8,9 With a VDE of 7.02 eV and its outstanding stability and weak coordination,20,21 B(CN)4–, known as tetracyanoborate or bison, is a hyperhalogen and has attracted both theoretical and experimental attention as a potential building block of halogen-free electrolytes. The anion shows a very high oxidative potential of 5.65 V (vs. Li+/Li0), 20 which is on par with that of the PF6– anion, currently used in the electrolyte of today’s lithium ion batteries. Its dissociation energy with lithium ion (Li+) is among the lowest compared to regular anions.20,22 The interest in halogen free bison B(CN)4– is further enhanced since currently used halogen-based electrolytes are notorious for their poor safety.8,9,23 However, it has been found that the LiBison salt shows very poor solubility in ionic liquids (IL) when forming electrolytes,

22

thus, hindering its applications in

Li-ion batteries. One explanation is that LiBsion tends to form large aggregates in ILs with four bison-coordinated Li[B(CN)4]43− ion as the building block.22 As shown in Figure 1a, the particular configuration of Li[B(CN)4]43− manifests due to Li+ and bison pair, LiB(CN)4, having the monodentate form shown in Figure.1a. Note that this from is energetically favored in various solvents.22 The basis of forming such monodentate configuration is the strong electron-drawing ability of the cyanide (CN−) group, caused by the concentrated charge on the nitrogen (as demonstrated by the charge distribution in Figure 1a). This monodentate configuration is further 2


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strengthened in solvents, due to the charge separation in an electrostatic environment. Therefore, one possible way to avoid the formation of the four-coordinated Li+ complex and to improve the solubility of the Li salt is to weaken the bonding between Li+ and the bison anion, while reducing the bonding polarizability. This would allow the monodentate configuration not to be favored. It has been reported recently that, by replacing the CN− with much less toxic thiocyanide SCN− inside a thermally stable dianion B12(CN)122–, one can obtain a new ultra-stable weakly coordinating dianion B12(SCN)122–, which shows significantly smaller dissociation energy with Li+ inside solutions.23 Inspired by such a result, in this paper, we use thiocyanide SCN− to replace the CN− group of the bison anion, with the aim of increasing the solubility of the Li salt. Properties of the as-made anion B(SCN)4− are discussed in the context of electrolyte application.

Figure 1. NBO charge distributions of (a) the ground states of the four-bison-coordinated Li+ complex Li[B(CN)4]43− together with the monodentate configuration of LiB(CN)4 in vacuum. (b) Similar results for the ground states of Li[B(SCN)4]43− together with the bi-dentate configuration of LiB(SCN)4 in vacuum. (c) The crystal structure of LiB(CN)4 showing the local environment of Li atom. The blue dotted line shows the cubic unit cell. Li is in green, B in orange, C in black and N in blue.

3



METHOD The geometries and total energies of Li[B(SCN)4]n(n-1)− and Li[B(CN)4]n(n-1)− (n = 1 ~ 4 ) are calculated using density functional theory with B3LYP hybrid exchange-correlation functional.24,25 We used the Gaussian 03 code26 with 6-311++G(d,p) basis set for all atoms.27 This functional and basis set have been proven to be capable of providing reliable results in previous works.22 In all cases, structures are fully optimized without any imaginary frequency. The convergence in the total energy and force are set at 1×10−6 eV and 1×10−2 eV/Å, respectively. The natural population analysis (NPA) of charges on each atom is carried out using the natural bond orbital (NBO) method. To explore solvation effects, the optimized vacuum structures are used as input using C-PCM (the polarizable continuum model implementation of the conductor-like screening model).28-30 Four solvents with different dielectric permittivity are selected to study the effect of different electrostatic environment on Li+Bison and Li+B(SCN)4, including cyclohexane (CHN, ε = 2.0), tetrahydrofuran (THF, ε = 7.4), acetonitrile (ACN, ε =36) and water (H2O, ε = 78). Dissociation energies, Ed are calculated as Ed = E(Li+) + E(Y−) − E(LiY), where Y = Bison and B(SCN)4.

RESULTS AND DISCUSSION

The ground state of B(SCN)4− We begin with the equilibrium structure of B(SCN)4− given in Figure 2. Here the boron is attached to nitrogen, rather than to sulfur. The optimized geometry of the anion is in good agreement with the result obtained by using the MP2 method.31 The symmetry operations of the structure belong to the Td group, as that of the bison anion. According to group theory, in order to form the directed valence bonding with the four SCN units stretching to the tetrahedral corners, the boron atom must be promoted from its ground state of 1s22s22p1 to an excited state involving the d state. The 4


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electronic excitation energy is over- compensated by the attractive bonding energy between boron and SCN−. Given that CN is more electronegative than SCN, the attractive bonding between boron and CN is expected to be stronger than that between boron and SCN. Therefore, B(SCN)4− is expected to be thermodynamically less stable than B(CN)4−. Indeed, the calculated VDE of B(SCN)4− is 5.24 eV which is lower than that of B(CN)4−. The VDE value computed here is consistent with the previous calculated results at the MP2/6-311+G* level of theory, where the VDE of negatively charged species M(SCN)4− (M = Li, Na, Be, Mg, Ca, B, Al) containing the NCS ligands range from 5.0 to 6.3 eV.31

Interaction of Li+ with B(CN)4− and B(SCN)4− As mentioned in the introduction, the precipitation and low solubility of the LiBison salt in ILs may be due to its tendency to form a complex with one Li+ ion coordinated with four B(CN)4− ions, as shown in Figure 1a. This configuration is made possible because the ground state of Li-bison ion pair adopts the monodentate configuration in solutions.19 As shown in Figure 2, in the vacuum, the monodentate configuration of Li-bison ion pair is 0.34 eV higher in energy than the bidentate configuration. The monodentate configuration in the solutions helps to minimize the repulsion between charge-concentrated nitrogen atoms when forming the tetrahedral Li[B(CN)4]43− complex. Note that this geometry is consistent with the local environment of lithium inside the crystal structure of the LiBison salt as shown in Figure 1c.18 On the other hand, for the binding of B(SCN)4− and Li ion pair, the bidentate configuration is found to be the ground state in both vacuum and in solutions, as shown in Figure 2. The monodentate configuration of Li[B(SCN)4] is unstable. It is a transition state to its bidentate configuration, as shown by the imaginary vibrational mode in Figure 2. The tridentate configuration is 0.09 eV and 0.34 eV higher in energy in the vacuum and in the solution (ɛ = 2), respectively. The preference of Li+-[B(SCN)4]− ion pair to have a bidentate instead of monodentate configuration in solutions is likely to prevent the formation of tetrahedral complex 5



between Li+ and B(SCN)4−.

Figure 2. Optimized structures of B(SCN)4−, B(CN)4− and their ion pairs with Li. ‘m’, ‘b’, and ‘t’ standing for monodentate, bidentate and tridentate, respectively. Also shown is the relative energetics (in eV) of different configurations of the ion pairs. In each case, the number in the first row corresponds to the value in vacuum and the number in the second row corresponds to the value in the solution with ε = 2.0. The tridentate configuration of LiBison in the solution is found to be unstable as indicated by ‘*’. For the case of LiB(SCN)4, its monodentate configuration is dynamically unstable, which serves as a transition state (TS) to the bidentate configuration. Boron is in pink, nitrogen in blue, carbon in gray, sulfur in yellow and lithium in violet.

The calculated dissociation energy of the Li+-B(SCN)4− ion pair is lower than that of LiBison as shown in Figure 3. When solvation effects are considered for four solvents, including cyclohexane (CHN, ε = 2.0), tetrahydrofuran (THF, ε = 7.4), acetonitrile (ACN, ε = 36) and water (H2O, ε = 78), the dissociation energy of Li[B(SCN)4] is even lower than that of LiBison. In fact, the ion pair becomes 6


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intrinsically unstable for acetonitrile (ε = 36) and water (ε =78), suggesting complete separation of the two ions. To further see if there is any difference when using other alkali elements, we repeated the above calculations for both Na+B(SCN)4− and K+B(SCN)4− ion pairs and compared those to NaBison and KBison, respectively. The results are similar as that in Li+-B(SCN)4−, both in the case of vacuum as well as in the cases with various solutions. The dissociation energies of Na+B(SCN)4− and K+B(SCN)4− ion pairs are less than those of NaBison and KBison.

To see if the lithium salt could be used in electrolytes of lithium/sulfur (Li/S) batteries, we calculated the dissociation energy of Li+-B(SCN)4− in solvent ether (ε = 4.3), since the suitable solvents for Li/S cell electrolytes are currently limited within linear or cyclic ethers.32 We found that the calculated dissociation energy is as small as 0.73 eV, suggesting a high solubility of the ion pair in the ether solvent.

Figure 3. Comparison of the dissociation energy between ion pairs of Li+-B(CN)4− and Li+-B(SCN)4−, Na+-B(CN)4− and Na+-B(SCN)4−, as well as K+-B(CN)4− and K+-B(SCN)4− in vacuum and various solvents.

7



The equilibrium geometry of Li[B(SCN)4]43− complex is shown in Figure 4. Each B(SCN)4− unit binds to Li+ with a Li-S-CN angle of 114◦, making a compact configuration, where nitrogen atoms of the neighboring SCN are close to each other. Since the charges are concentrated on N, as shown by the NBO result in Figure 1b, such configuration generates greater repulsion between the anions, while sulfur, bearing little charge, can only provide limited energy gain from attraction to Li+, making such complex less stable. Indeed, the calculated dissociation energy of Li[B(SCN)4]43− → Li+ + 4·B(SCN)4− is about 1 eV smaller than the dissociation energy of Li[B(CN)4]43− → Li+ + 4·B(CN)4−. The energy needed for three B(SCN)4− and one Li[B(SCN)4] to form the four-coordinated Li+ complex Li[B(SCN)4]43− is almost twice the energy of that in the case of Li[B(CN)4]43−, as shown in Figure 5. This suggests that it is energetically not favorable for Li[B(SCN)4]43− to form the four-coordinated Li+ complex.

Figure 4. Optimized structures of Li[B(CN)4]n(n-1)− and Li[B(SCN)4]n(n-1)− in vacuum for n = 8


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2~4. Boron is in pink, nitrogen in blue, carbon in gray, sulfur in yellow and lithium in violet.

The optimized structures of Li[B(SCN)4]n(n-1)− and Li[B(CN)4]n(n-1)− (n = 2 ~ 4 ) in the vacuum are given in Figure 4. Li[B(SCN)4]43− is a metastable state which is less stable against the reaction channel Li[B(SCN)4]43− → Li[B(SCN)4]32− + B(SCN)4−. Li[B(SCN)4]32− is also metastable, which can undergo the reaction Li[B(SCN)4]32− → Li[B(SCN)4]2− + B(SCN)4−. We find that further reaction of Li[B(SCN)4]2− → Li[B(SCN)4] + B(SCN)4− will cost at least an energy of 1.2 eV, suggesting that Li[B(SCN)4]2− is the preferred composition. The energetics of various reaction channels are given in Figure 5. As shown in Figure 4, the Li+ and B(SCN)4− ion pair forms the same bidentate binding block in Li[B(SCN)4]2−, which is consistent with what we found in the single ion pair of Li[B(SCN)4].

Figure 5. Relative energetics of different fragmentation channels of (a) Li[B(SCN)4]43− and (b) Li[B(SCN)4]43−.

According to the above discussion, we conclude that B(SCN)4− is unlikely to form multiple-coordinated Li+ complex due to the special bonding features of the Li and B(SCN)4− ion pair; the absence of the monodentate configuration and large 9



repulsion generated by charge-concentration on nitrogen atoms close to each other. Although B(SCN)4− is unlikely to form aggregated complex with Li+ as discussed above, we find that the oxidative stability of the anion is lower than that of Bison. The calculated HOMO-LUMO gap of B(SCN)4− is 6.05 eV against 8.90 eV of Bison. The oxidation potential (vs. Li+/Li0) of B(SCN)4− is a moderate 3.87 eV compared to the high value > 5.50 eV of Bison. Therefore, B(SCN)4− may not be suitable for high-voltage applications.

Figure 6. Optimized structures of the mixed anions B[(CN)x(SCN)4-x] –. Boron is in pink, nitrogen in blue, carbon in gray and sulfur in yellow.

Anions with mixed CN and SCN ligands We reckon that mixed anions of B[(CN)x(SCN)4-x]– (x = 1 ~ 3) may have better oxidative stability, while at least partially maintaining the advantage of B(SCN)4−. Figure 6 shows the optimized structures of these mixed anions. Indeed, it is found that they all have higher oxidative stability than that of B(SCN)4–; the oxidation potentials (vs. Li+/Li0) of B[(CN)x(SCN)4-x] – (x = 1 ~ 3) are 3.93 eV, 4.17 eV and 4.65 eV, respectively.

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CONCLUSION By replacing cyanide with much less toxic thiocyanide, SCN, in the bison anion B(CN)4–, a weakly-coordinating anion, B(SCN)4–, is introduced as a new member of the tetracyanoborate family. Compared to the bison anion, B(SCN)4– shows lower dissociation energy with Li+ as well as other alkali cations (Na+ and K+) in both vacuum and solutions. The four-coordinated Li+ complex Li[B(SCN)4]43− is unlikely to form due to the instability of the monodentate configuration of the B(SCN)4– and Li+ ion pair. All these results suggest that that lithium salts with B(SCN)4– would have better solubility than that of LiBison in the IL electrolytes. We also find that mixing the cyanide and thiocyanide groups in the anion can significantly enhance the moderate oxidative stability of B(SCN)4–.

ACKNOWLEDGEMENTS This work was supported by the National Natural Science Foundation of China (No.11504301), Fundamental Research Funds for the Central Universities (XDJK2015C045 and SWU114088) and the U. S. Department of Energy, Office of Basic Energy Science, Division of Materials Sciences and Engineering under Award #DE-FG02-96ER45579. Resources of the National Energy Research Scientific Computing Center Supported by the Office of Science of the U. S. Department of Energy under Contract No. DE-AC02-05CH11231 is also acknowledged.

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