1. Organic solvents
Lithium secondary batteries have higher operating voltages, so organic solvents are used instead of aqueous solutions as electrolytes. The biggest drawback of organic solvents is their low dielectric constant. The electrolyte needs to have high ionic conductivity to dissolve the lithium salt and act as a polar aprotic solution to prevent reaction with lithium. Figure 1 summarizes the representative physical and chemical characteristics of lithium battery electrolyte. The dielectric constant of the solvent affects the dissociation and association of ions in the lithium salt. The higher the dielectric constant, the more conducive to rapid separation, because the dielectric constant is inversely proportional to the Coulomb force between the cation and anion of the lithium salt.
Generally speaking, the dielectric constant of the medium solvent needs to be greater than 20, because when the dielectric constant is small, the dissociation of the lithium salt is difficult to proceed. According to Stokes’ theorem, the movement of ions in a liquid electrolyte is inversely proportional to the viscosity of the solvent. Therefore, the viscosity of the solvent should be 1 cP or less. The donor number and acceptor number represent the protic and electrophilic properties of the solvent, respectively. These values provide information about the interaction of cations and anions or the strength of the solvation of the lithium salt. The dissociation of lithium salt increases as the number of acceptors increases. The working temperature is affected by the melting point and boiling point of the organic solvent, which requires the solvent to maintain a liquid state at room temperature and be able to dissolve lithium at a temperature as low as -20 ℃. The organic solvent should have a high boiling point and a low vapor pressure. At the same time, the electrolyte must have a higher dielectric constant and lower viscosity to obtain a higher ionic conductivity. A high dielectric constant will lead to an increase in polarization and viscosity. These can be achieved by mixing a solvent with a high dielectric constant and a solvent with a low viscosity. For example, cyclic carbonates such as ethylene carbonate (EC) and propylene carbonate (PC) have high dielectric constant and high viscosity, and their solvent molecules have strong interactions. The dissociation of lithium salt may occur in EC but EC cannot be used alone because its melting point is too high. In contrast, linear carbonates such as dimethyl carbonate (DMC) and diethyl carbonate (DEC) have low dielectric constant and low viscosity. Therefore, the cyclic carbonate and the linear carbonate are used in combination to obtain the characteristics required as an organic solvent for a lithium secondary battery. Figure 2 shows the adoption of 1M LiPF. Dissolve in an organic solvent to prepare the ionic conductivity of the liquid electrolyte. It can be seen from Figure 2 that the mixed solvent has a higher ionic conductivity.
2. Lithium salt
Figure 3 shows the physical and chemical properties of lithium salts commonly used in lithium secondary batteries. Anions with a larger ionic radius are advantageous because lithium salts with delocalized anions are easier to dissociate. Generally speaking, the dissociation of potassium salt proceeds in the following order.
Li(CF3SO2)2N＞LiAsF6＞LiPF6＞LiClO4＞LiBF4＞LiCF3SO3
On the other hand, an increase in ion radius leads to a decrease in anion mobility. As shown in equation (3.16), ion mobility conforms to Stokes’s theorem, which can also be described in terms of diffusion coefficient.
μ0 = λ0/(zF)=ze/(6πη0r)=zFD/(RT)
In the formula, λo, z, F, e, r, 7o, R and T represent the limit molar conductivity, charge quantity, Faraday constant, element charge, ion radius, viscosity, gas constant and absolute temperature, respectively. As mentioned above, the size of the anion is an important factor in determining the properties of the lithium salt. Figure 4 shows the ion radius of the lithium salt drawn based on the space filling model and van der Waals effect.
As a lithium salt, LiClO4 is commonly used in lithium primary batteries. However, due to the high oxidation environment generated during the charging process, safety problems are prone to occur, so it has not been applied to lithium secondary batteries. Other lithium salts such as LiBF4 and LiPF6 containing fluorinated Lewis acid are often used in lithium secondary batteries because of their good workability and chemical stability. Other salts under study include inorganic lithium salts, organic sulfonates and imide salts. The ionic conductivity of LiBF4 electrolyte is lower than that of electrolytes containing LiClO4 or LiPF6, which will affect the high rate performance of the battery. In contrast, LiPF6 electrolyte has high ionic conductivity, but poor thermal stability and no side reactions on the electrode. The performance of the battery will further deteriorate with the occurrence of side reactions. When the electrolyte is exposed to moisture, it will decompose to produce HF, which causes the electrolyte to decompose. Figure 5 compares various characteristics of representative lithium salts.
At the same time, LiRfSO3 has not yet been commercialized, its solubility is low, and the lithium ion conductivity is low. Lithium sulfonimide (Li[RfSO2]2N) has good chemical stability, but its ability to resist oxidation in the positive electrode is weak, and it is easy to corrode the aluminum current collector, so it cannot be used in actual batteries.
3. Application of molecular orbital theory in solvent design
The molecular orbital equation uses the Schrodinger equation Hψ=Eψ (H: Hamiltonian; E: the sum of electric potential and dynamic energy) to describe the wave-like behavior of electrons. The wave function ψ is the possibility of finding an electron in a molecule, and its actual possibility is obtained by [ψ]². When the molecular orbital function is similar to the atomic wave function, electrons are scattered in molecules instead of atoms. Electrons begin to occupy the wave function from a very low energy level. Here, the ones occupying the highest energy molecular orbital and the lowest energy molecular orbital are called HOMO and LUMO, respectively. Figure 6 illustrates the different energy levels of HOMO and LUMO.
The energy levels of different substances are unique and change with the type of solvent in the electrolyte. The potential window of a given electrolyte can be calculated according to the HOMO-LUMO theory.
Having a high HOMO energy level can promote the oxidation reaction by donating electrons, while having a low LUMO energy level can promote the reduction reaction by accepting electrons. Solvents with low HOMO and high LUMO are suitable for use in electrolytes.
Figure 7 shows the relationship between solvents with different HOMO and LUMO energy levels
