How batteries work Chapter 8 Liquid electrolytes

Chapter 8

Liquid electrolytes: formulation, salt dissociation and Li+ solvation

What is the liquid inside, and why these molecules?

Where we are. Chapter 7 put our ion in its two hosts and sent it from one to the other. Between them it crosses the electrolyte, the part of the cell that must carry ions and stop electrons. What is that medium, and how does an ion get through it? We start from nothing: what the liquid in a lithium-ion cell is made of, how its molecules sit around each ion, how an ion moves through it, what the electric field inside it looks like, and how the arrangement changes, layer by layer, as we approach an electrode. Then we compare the liquid with polymers and ceramics, and see why the anion matters as much as the lithium.

Figure 8.1 Three views of four species: a turning ball-and-stick model, a computed map of electron density, and a computed map of the electric potential around each one, red where negative and blue where positive. The ball-and-stick view shows one of four species, chosen with buttons: ethylene carbonate, a five-membered ring with a carbonyl oxygen sticking out; dimethyl carbonate, the same carbonate group with a methyl group on each side; the PF6 anion, a phosphorus atom with six fluorine atoms around it; and the lithium ion. A fifth view shows all four side by side at the same scale. A key gives the colour of each element; the carbonyl oxygen is ringed.
The cast: two solvents, one anion, one lithium ion. Try this: drag the molecule round until the ringed oxygen points straight at you.
SourceEC cyclic and DMC linear carbonates after Ue et al. 2014, section 2.1 and Table 2.1; the PF6− structure as drawn in Henderson 2014, Fig. 1.3; binding of Li+ through the lone pairs of the solvent’s donor atom after Henderson 2014, p. 9. Geometries, electron densities and electric potentials computed for this page with density-functional theory (B3LYP functional, def2-SVP basis set, PySCF; geometries optimized; one isolated molecule in vacuum, no solvent around it). Density maps are slices through each molecule’s plane on a logarithmic scale; potential maps show the potential in that plane outside the 0.001 electron per bohr3 surface (white line), clipped at ±2 V. Atoms in the ball-and-stick view are drawn smaller than their real size so the bonds show.

Let us begin with the cast. Everything in this module is made of four kinds of object, so we meet them first. A molecule is a group of atoms held together by chemical bonds; the figure draws each atom as a ball and each bond as a stick. The two solvents of a lithium-ion cell are both carbonates, built round a carbon atom bonded to three oxygen atoms. In ethylene carbonate, EC, the carbonate closes into a ring; in dimethyl carbonate, DMC, it stays an open chain: EC is a cyclic carbonate and DMC a linear one.1

In both, one oxygen is double-bonded to the carbon and sticks out of the molecule: the carbonyl oxygen. A lithium ion is held by a solvent through the spare pairs of electrons on such a donor atom.2 The salt brings the other two objects: the lithium ion, Li+, a lithium atom that has lost one electron, and the hexafluorophosphate anion, PF6−, six fluorine atoms around a phosphorus atom, carrying one negative charge.2

Electron clouds and charge

A molecule is not a set of hard balls. Its electrons form a cloud, densest at the atoms and thinning out over a fraction of a nanometre, and the figure shows that cloud as calculated for this page. The calculation also shows where each molecule is slightly negative and where slightly positive. EC has a negative end, at the carbonyl oxygen, and a positive end, at its CH2 groups: a strongly polar molecule, with a computed dipole moment of about 5 debye. In DMC the two halves almost cancel, about 0.3 debye. The same contrast shows up in the two liquids’ dielectric constants, 90 for EC against 3.1 for DMC, in the next figure, and simulations likewise treat EC as the polar solvent and DMC as the less polar one.1,3

Four objects to recognize: EC, a ring; DMC, a chain; PF6−, a ball of fluorine; and Li+, the smallest of all.

Figure 8.2 Three panels. Left: a beaker holding the recipe, 1 mol of LiPF6 per litre of a 3 to 7 mixture of EC and DMC. Middle: a bar chart of how many molecules there are for each lithium ion, about 4.5 EC, 8.2 DMC, one lithium ion and one PF6 anion. Right: a window a few nanometres across showing those molecules mixed at random; buttons switch to EC alone, a solid, or DMC alone, where the ions stay in pairs.
Li+PF6−ECDMCcarbonyl oxygen
What the liquid is made of, counted molecule by molecule. Try this: switch to DMC alone and watch where the anions go.
SourceRecipe after Ue et al. 2014, section 2.1; densities, molar masses, relative permittivities, viscosities and melting points from their Table 2.1 (EC values at 40 °C); the 3 : 7 ratio is one illustrative choice inside their advice to keep EC below about 30 to 50 % by volume. The counts are computed here from those numbers, neglecting the volume of the salt. Ion pairing in DMC after Borodin 2014, pp. 383 and 384. The window is schematic: positions are random and the glyphs are not to scale.

The recipe

Now we put the cast together. An electrolyte has one job that defines it: to provide purely ionic conduction between the two electrodes.4 It is a solution: a salt dissolved in a liquid, the solvent. Most commercial lithium-ion cells use roughly 1 mol of the salt lithium hexafluorophosphate, LiPF6, per litre of a mixture of carbonate solvents: a cyclic one such as ethylene carbonate (EC) or propylene carbonate (PC), and a linear one such as dimethyl carbonate (DMC), ethyl methyl carbonate (EMC) or diethyl carbonate (DEC).1

Water, the obvious solvent, is out: it would be split at the voltage of a lithium cell (chapter 3).5 Water-based electrolytes conduct best, about 1 S/cm; the organic liquids reach 10−2 to 10−3 S/cm, because they dissolve salts less well and let the ions pair up.4

Why two solvents

Why two solvents? Because each does one job well and the other badly. A good solvent needs a high dielectric constant (relative permittivity), to pull the salt apart into separate ions, and a high fluidity, so that the ions can move.1 EC has a relative permittivity of 90 but is thick, 1.9 mPa s, and melts at about 36 °C, so on its own it is a solid at room temperature. DMC flows easily, 0.59 mPa s, but its relative permittivity is only 3.1.1

Mixed, they stay liquid far lower: a mixture of carbonates has a eutectic, the composition with the lowest melting point, at −8.6 °C for EC with DMC and −64.7 °C for EC with DMC and EMC (the solvents alone, without salt), and blends near the eutectic are preferred to work in the cold. EC is usually kept below 50 % by volume, preferably 30 %, because more of it comes out as a solid at low temperature.1

A pinch of additives

Real electrolytes also carry minute amounts of additives, each with a job: some decompose on purpose to build a thin, controlled SEI on the graphite (chapter 10), others protect the positive electrode or stop overcharge (chapter 17); no single additive can do everything, and finding the right combination is now central to electrolyte work.6

Why LiPF6

Why LiPF6? It is used almost exclusively in commercial lithium-ion cells because it has shown the best balance of the properties a salt needs: its solutions are among the most conductive known, it forms a stable surface on the aluminium current collector at high potential, and with carbonate solvents it forms a stable SEI on graphite (chapter 10). Its weakness is its P–F bond, which is easily broken: the salt reacts readily with water and is not very stable to heat, and hydrofluoric acid, HF, in LiPF6 electrolytes is one of the principal concerns about using it.2 It became the standard only when carbon replaced lithium metal and the solvents were optimized for the lithium-ion cells Sony introduced in 1991.2

Worked example. Let us count the crowd: 1 mol of LiPF6 dissolved in 1 litre of EC : DMC 3 : 7 by volume (the salt’s own volume neglected). EC: 300 cm3 × 1.32 g/cm3 ÷ 88.1 g/mol = 4.5 mol DMC: 700 cm3 × 1.06 g/cm3 ÷ 90.1 g/mol = 8.2 mol so for each Li+ and each PF6−: about 13 solvent molecules each Li+ has a cube of liquid about (1 L ÷ 6.02 × 1023)^1/3 = 1.2 nm on a side to itself, on average A lithium ion is never more than a few molecules from the next one, and most of what surrounds it is solvent.1,5

The liquid in a lithium-ion cell is mostly solvent: a dozen or so carbonate molecules for every lithium ion, two kinds of solvent because neither does both jobs.

Figure 8.3 Two views. First, a table of six lithium salts, LiPF6, LiBF4, LiClO4, LiAsF6, LiTFSI and lithium triflate, with a bar for each one's conductivity at 1 molar and a symbol for how it does on resisting water, stability to heat, protecting the aluminium current collector, safety, and commercial use. Second, bar charts of five carbonate solvents, EC, PC, DMC, EMC and DEC, for a chosen property: permittivity, viscosity, melting point, boiling point, flash point or oxidation limit.
good, or a cyclic solventpartly, or a linear solventpoor
The ingredients of commercial electrolytes, side by side. Try this: in Salts, tap the aluminium symbol for LiTFSI; then in Solvents, compare the flash points of EC and DMC.
SourceSalts after Henderson 2014, section 1.3 (pp. 7 to 9); conductivities of 1 M solutions in PC:DME (50:50 by volume) at 25 °C from his Table 1.1, and the aluminium repassivation potentials from Table 1.7. The good, partly and poor symbols are this page’s summary of his wording, which the readout quotes for every cell; “not rated” means the passage gives no verdict. Solvents from Ue et al. 2014, Table 2.1: EC permittivity and viscosity at 40 °C; oxidation limits on glassy carbon with tetraalkylammonium salts, not with lithium salts.

Choosing a salt

We have taken one salt and two solvents as given. Why these, and not others? Let us line up the alternatives, starting with the salt. A salt for a lithium-ion electrolyte has a list of jobs. It must give a conductive solution, stay stable within the cell’s voltage window and at high temperature, help form a good SEI, protect the aluminium current collector of the positive electrode above about 3.6 V vs Li, and resist water. It should also be cheap and not toxic. Failing any one of these rules a salt out of practical use, and every one of them depends on the rest of the recipe.2

Put the candidates side by side and no salt wins everywhere. LiClO4 conducts well and is stable, but its anion is a strong oxidant, which makes the salt a potential explosive. LiAsF6 behaves much like LiPF6, but reduction could turn its arsenic into highly toxic forms. LiBF4 resists water and heat better than LiPF6 but conducts significantly less.2

LiTFSI and lithium triflate shrug off water and heat, but in ordinary dilute solutions they corrode the aluminium foil at high potential. In PC:DME the aluminium repassivation potential, above which corrosion pits keep growing, is 3.7 V vs Li with LiTFSI, against above 5 V with LiPF6.2 So LiPF6, with its known weaknesses, keeps its place in the recipe of figure 8.2.2

Choosing the solvents

The solvents tell the same story. The five carbonates of commercial electrolytes split into two families.1 The cyclic ones, EC and PC, have relative permittivities of 65 to 90 and flash points above 130 °C, but they are viscous, and EC is a solid below 36 °C. The linear ones, DMC, EMC and DEC, flow two and a half to four times more easily and, apart from DMC, stay liquid far below 0 °C. But their permittivities are only about 3, and they flash at 17 to 25 °C, near room temperature.1

All five resist oxidation on glassy carbon to above 6 V vs Li. A window measured on inert carbon, however, is not a clear guide to stability against real electrode materials.1,2

Every salt and every solvent fails at something; a commercial electrolyte is the mixture whose weaknesses matter least.

Beyond carbonates: ionic liquids

One family of liquids avoids flammable solvents altogether. Ionic liquids, salts that are liquid at room temperature, are non-flammable and thermally stable and offer electrochemical windows of 5 to 6 V; their cations are typically imidazolium, quaternary ammonium, pyrrolidinium or piperidinium, their anions TFSI, BF4− or PF6−. Their weakness is transport: they are more viscous than organic liquids and have lower lithium transference numbers, which limits their rate.7

Figure 8.4 Left: a flat sketch of solid LiPF6, lithium ions and PF6 anions alternating in rows. Right: liquid solvent. Again and again, four solvent molecules close in on an exposed lithium ion at the edge of the solid, turn their oxygens to it and carry it away into the liquid; the anion beside it drifts off alone.
Li+PF6−ECDMCcarbonyl oxygen
Dissolving the salt: the solvent takes the lithium ions one by one. Try this: follow one lithium ion from the edge of the solid until it is out in the liquid.
SourceDissolution through the solvation of Li+ by the lone pairs of the solvent, with the anions in general left unsolvated because aprotic solvents have no acidic protons, after Henderson 2014, p. 9; solvent–anion complexes such as EC with PF6− (Borodin 2014, p. 376) show that solvent molecules do come close to the anion; the faint, circling neighbours are a drawing device for that weak, short-lived contact, oriented by the computed potential of figure 8.1; attraction of opposite charges, chapter 0. The solid is a flat sketch, not the crystal structure of LiPF6, and the sequence is slowed by many orders of magnitude.

We have the recipe; now let us watch the salt go into the liquid. What does it mean for a salt to dissolve? Solid LiPF6 is made of Li+ cations and PF6− anions, which attract each other because their charges are opposite (chapter 0). In a carbonate solvent the Li+ ions are pulled out of it by the solvent molecules: each solvent molecule has an oxygen atom with spare pairs of electrons (in a carbonate, the carbonyl oxygen), and it binds to the lithium through them. The anions are, in general, left bare, because these solvents have no acidic hydrogen to hold them: in water an anion is held by hydrogen bonds, in a carbonate it is not. The anions can bind to Li+ in the same way, so solvent and anion compete for the places around each lithium ion, and that competition decides what is actually present in the liquid.2

And the anion?

So does a PF6− anion get a shell of its own? Not in the way Li+ does. Solvent molecules do come close to it, and simulations treat a PF6− next to an EC molecule as a complex, but nothing binds there the way the oxygens bind to lithium, and the neighbours keep changing. In the figure they are drawn faint, with their slightly positive CH ends (blue in figure 8.1) towards the anion.2,3

Dissolving is a tug of war for the lithium ion, and in a good solvent the solvent wins.

Figure 8.5 A turning three-dimensional model of a lithium ion surrounded by four solvent molecules at the corners of a tetrahedron, each pointing its carbonyl oxygen at the ion at a distance of 0.204 nanometres, with a PF6 anion outside the shell. Buttons switch to a contact ion pair, where the anion takes one place in the shell, and an aggregate, where one anion is shared by two lithium ions. A side panel compares how long partners stay: an EC molecule about half a nanosecond, an anion 2 to 3 nanoseconds.
One lithium ion and its shell, in three dimensions. Try this: drag the shell round until you can see all four molecules, then watch one pull away and come back.
SourceCoordination through the lone pairs of the solvent, bare anions and the competition for the shell after Henderson 2014, pp. 9 to 13; four solvent molecules in a tetrahedron, 4.2 to 4.4 partners in EC : DMC and EC : PC with LiPF6 from molecular dynamics, 4.5 PC molecules at 2.04 Å from neutron diffraction, the residence times and the names SSIP, CIP and aggregate after Borodin 2014, p. 382; bound through a single fluorine (monodentate) as the simulations find most stable in the liquid, also p. 382. Atom positions other than the Li–O distance are approximate. The exchange is shown slowed by about ten billion times.

The shell

Let us follow one lithium ion out of the crystal and look at what it has gathered. How many solvent molecules fit round it? Few. Simulations find about four solvent molecules bound strongly around a lithium ion, arranged as a tetrahedron, and about 4.2 to 4.4 partners, solvent and anions together, in EC : DMC and EC : PC with LiPF6; neutron diffraction on LiPF6 in PC finds 4.5 PC molecules on average, each oxygen 0.204 nm from the lithium.3

This group is the ion’s first solvation shell. In the liquid, simulations find EC and DMC in it about equally often, even though an isolated lithium ion in the gas phase prefers EC.3 The shell is not a fixed object: a given EC molecule stays in it for about half a nanosecond, an anion partner for 2 to 3 ns.3

Pairs and clusters

When an anion wins a place, the two ions form a contact ion pair; when solvent stays between them, they are a solvent-separated ion pair; larger clusters are aggregates.2,3 The anion matters: the tendency to pair up rises in the order LiAsF6 and LiPF6, then LiTFSI, LiClO4, LiBF4 and lithium triflate.2

The solvent matters too: in simulations of EC : DMC the fraction of free ions rises nearly linearly with the share of EC; and measured with LiTFSI, the conductivity reaches 62 % of what fully independent ions would give in PC, but only 11 % in DMC. Not every cluster is lost to conduction: charged clusters carry current too.3

A lithium ion never travels alone: it carries a shell of about four solvent molecules, and sometimes an anion.

You can now explain what the liquid in a lithium-ion cell is made of and why it mixes two solvents, what a solvation shell is and how ions pair up, how an ion wanders and drifts and how weak the field in the bulk is, how the liquid changes layer by layer towards an electrode, how a liquid, a polymer and a ceramic each carry a lithium ion, why dry-polymer cells run warm, what the transference number is and why a low one leaves a salt gradient, and why a solid electrolyte needs more than a high conductivity.

Our ion has now crossed the electrolyte. On the far side, it has to find a particle.

The mathematics of this chapter

Nsolvent/NLi = (1/c) Σ φkρk/Mk, d = (1/(cNA))1/3solvent molecules per lithium ion, for volume fractions φk of solvents of density ρk (g/L) and molar mass Mk (g/mol), salt concentration c (mol/L) (salt volume neglected); the mean spacing of the ions (figure 8.2)

E = i/κ, v̄ = i/(Fc)the field in the bulk liquid from Ohm’s law; the mean velocity of the lithium ions once all the current is carried by lithium (worked example of figure 9.1)

κ = F Σ|zj|ujCj, R = l/(κA)conductivity from the ions; the resistance of a layer of thickness l and area A (worked example)

tj = |zj|ujCj / Σ|zk|ukCk, t+ = DLi/(DLi + D−) (dilute)the transference number, and its dilute-solution form

Jj = −Dj ∇Cj − (zjF/RT) Dj Cj ∇φ, i = F Σ zjJjNernst-Planck, as in chapter 0: each ion’s flux by diffusion and by migration in the field; the current density from the fluxes

Δc = (1 − t+) i L / (F D)the steady-state salt difference across a layer of thickness L at current density i (figure 9.2)

References for this chapter

Numbers in the text refer to this list. All references of the course, with the chapters that use each one.

  1. 1 M. Ue, Y. Sasaki, Y. Tanaka, M. Morita, "Nonaqueous Electrolytes with Advances in Solvents", in T. R. Jow, K. Xu, O. Borodin, M. Ue (eds.), Electrolytes for Lithium and Lithium-Ion Batteries, Modern Aspects of Electrochemistry 58, Springer, New York, 2014, 93-165. doi:10.1007/978-1-4939-0302-3_2
  2. 2 W. A. Henderson, "Nonaqueous Electrolytes: Advances in Lithium Salts", in T. R. Jow, K. Xu, O. Borodin, M. Ue (eds.), Electrolytes for Lithium and Lithium-Ion Batteries, Modern Aspects of Electrochemistry 58, Springer, New York, 2014, 1-92. doi:10.1007/978-1-4939-0302-3_1
  3. 3 O. Borodin, "Molecular Modeling of Electrolytes", in T. R. Jow, K. Xu, O. Borodin, M. Ue (eds.), Electrolytes for Lithium and Lithium-Ion Batteries, Modern Aspects of Electrochemistry 58, Springer, New York, 2014, 371-401. doi:10.1007/978-1-4939-0302-3_8
  4. 4 M. Winter, R. J. Brodd, "What Are Batteries, Fuel Cells, and Supercapacitors?", Chemical Reviews 104, 2004, 4245-4269. doi:10.1021/cr020730k
  5. 5 A. J. Bard, L. R. Faulkner, H. S. White, "Electrochemical Methods: Fundamentals and Applications", Wiley, 3rd edition, 2022.
  6. 6 K. Abe, "Nonaqueous Electrolytes and Advances in Additives", in T. R. Jow, K. Xu, O. Borodin, M. Ue (eds.), Electrolytes for Lithium and Lithium-Ion Batteries, Modern Aspects of Electrochemistry 58, Springer, New York, 2014, 167-207. doi:10.1007/978-1-4939-0302-3_3
  7. 7 O. Faruque, M. J. Isalm, O. Faruq, M. M. Alam, "A comprehensive review of past and present developments of Li-ion batteries", Discover Sustainability 7, 2026, 881. doi:10.1007/s43621-026-02920-8
  8. 8 J.-M. Tarascon, M. Armand, "Issues and challenges facing rechargeable lithium batteries", Nature 414, 2001, 359-367. doi:10.1038/35104644
  9. 9 J. B. Goodenough, K.-S. Park, "The Li-Ion Rechargeable Battery: A Perspective", Journal of the American Chemical Society 135, 2013, 1167-1176. doi:10.1021/ja3091438
  10. 10 B. Boz, T. Dev, A. Salvadori, J. L. Schaefer, "Review: Electrolyte and Electrode Designs for Enhanced Ion Transport Properties to Enable High Performance Lithium Batteries", Journal of The Electrochemical Society 168, 2021, 090501. doi:10.1149/1945-7111/ac1cc3
  11. 11 D. Shanmukaraj, P. Ranque, H. Ben Youcef, T. Rojo, P. Poizot, S. Grugeon, S. Laruelle, D. Guyomard, "Review—Towards Efficient Energy Storage Materials: Lithium Intercalation/Organic Electrodes to Polymer Electrolytes—A Road Map (Tribute to Michel Armand)", Journal of The Electrochemical Society 167, 2020, 070530. doi:10.1149/1945-7111/ab787a
  12. 12 K. Brandt, "Historical development of secondary lithium batteries", Solid State Ionics 69, 1994, 173-183. doi:10.1016/0167-2738(94)90408-1
  13. 13 O. Borodin, X. Ren, J. Vatamanu, A. von Wald Cresce, J. Knap, K. Xu, "Modeling Insight into Battery Electrolyte Electrochemical Stability and Interfacial Structure", Accounts of Chemical Research 50, 2017, 2886-2894. doi:10.1021/acs.accounts.7b00486

Cite this page

Using this page in teaching, a thesis or an article? Cite it as follows; the date is the last update of the course.

Shyma Sajeevan, A. (2026). Liquid electrolytes: formulation, salt dissociation and Li⁺ solvation. How batteries work. https://asajeevan.github.io/batteries/liquid-electrolytes/

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