Lithium

Typical wastewaters

• lithium carbonate production (brine and trona, spodumene) Li⁺ in residual brine returned to the source water body (trona process); spodumene process wastewater limited on TSS and pH only
• battery manufacturing and recycling hydrometallurgy Li⁺ with sodium, potassium, boron and sulfate in spent brine, concentrator tailings water and recycling liquors no effluent concentrations were read; the USGS flags battery use and disposal as a future source
• municipal sewage (pharmaceutical lithium, swimming pool sanitiser) Li⁺, conservative; lithium salts used as a pharmaceutical and lithium hypochlorite as a pool sanitiser are the EPA's named anthropogenic sources no sewage concentration was read

1 · Identity

Symbol, number

Li, 3

Oxidation states in water

+1 only, as the small, strongly hydrated Li⁺ ion. It forms no hydrolysis products, no insoluble hydroxide and no significant complexes at natural concentrations, so it behaves as a conservative cation like sodium (book entry: its salts are soluble and it moves freely once in water).

Note

The element entry carries the metal's reaction with water, the minerals, the brines and the 2024 production figures. This chapter is about lithium as a dissolved trace cation and as an industrial effluent.

2 · Occurrence in water

Natural sources

Weathering of lithium bearing silicates and evaporite minerals; concentrations are highest in arid regions and in older groundwater, particularly in unconsolidated clastic aquifers and sandstones, and lowest in carbonate rock aquifers; extensive evaporation, mineral dissolution, cation exchange and mixing with geothermal waters or brines account for the observed concentrations (USGS 2021). Geothermal and oilfield brines and the continental salt flat brines of Argentina, Chile and China are the concentrated natural reservoirs (book entry).

Anthropogenic sources

Brine evaporation ponds and spent brine, spodumene concentrator and lithium carbonate refinery effluent, battery manufacturing and recycling hydrometallurgy, lithium salts in swimming pool sanitisers and as a pharmaceutical (EPA); the USGS notes that anthropogenic sources may become important because of the rapid increase in battery use and disposal.

3 · Speciation

Lithium in water is Li⁺ at every pH and redox condition met in treatment. It does not hydrolyse, does not precipitate as hydroxide or carbonate at trace concentrations, and is not removed by heating, boiling or disinfection (EPA fact sheet). Only in evaporating brines does it reach the solubility of lithium carbonate, which is how it is won from salt flats.

Solubility

Lithium chloride, sulfate, nitrate and hydroxide are freely soluble; lithium carbonate is the least soluble common salt and the one precipitated in production. No solubility figure is quoted because none was read this session.

Hydrolysis

Negligible; Li⁺ is not an acid in water at any concentration met in treatment.

Complexation

Weak ion pairing with sulfate and carbonate in brines only; not significant in fresh water (general chemistry, not quantified in the sources read).

Precipitates

Li₂CO₃ from concentrated brine with sodium carbonate; lithium phosphate and lithium aluminate in recovery processes (not read this session, named only). None in natural water.

4 · Role in treatment

5 · Removal and control

6 · Analytics

Sampling pitfalls

Acidify as for any metal; lithium is conservative and stable in the bottle. Contamination comes from lithium grease, some laboratory glassware and lithium heparin blood tubes if a clinical lab is shared (general practice). The health screening level is close to the reporting limit, so method sensitivity decides what the data can say.

7 · Regulatory limits

Limits change, and many are set locally. Treat these as the published values to start from, not as your compliance target: check the standard in force at your site and the numbers written into your own permit.

8 · Health and environmental effects

Toxicity

No recommended dietary allowance and differing views on whether lithium is a micronutrient (EPA). Therapeutic doses of 600 to 1200 mg/day of lithium compound treat bipolar disorder and carry renal effects (dilute urine, thirst, kidney disease at higher long term doses), neurological effects (lethargy, tremor, cognitive impairment) and thyroid and parathyroid effects; the EPA provisional reference dose of 2 µg/kg per day is a LOAEL from patients divided by 1000, so risk between 10 µg/L and therapeutic levels cannot be confidently estimated (EPA fact sheet).

Bioaccumulation

Not addressed in the sources read; lithium is present in cereal grains, leafy vegetables and root crops (EPA).

Ecotoxicity

No US EPA aquatic life criterion; lithium does not appear in the national recommended criteria table.

Flags

• The WHO row is an absence: no fact sheet or guideline was found, but the full GDWQ table of guideline values was not re-read this session.
• US occurrence is a 1991 to 2018 USGS compilation of untreated groundwater; the UCMR 5 national dataset (August 2026) was not read.
• The reverse osmosis and tracer statements are general practice, not from a source read.
• No effluent concentration for brine, concentrator or battery plants was read; the ledger's mining chapter, not this entry, is where they should sit.
• Abu Dhabi sewer value 2.5 mg/L; marine outfall has no lithium row; other GCC states not read.

Gaps

• No surface water, sewage or industrial effluent concentrations were read.
• No solubility of lithium carbonate and no ion exchange selectivity data were read.
• No removal percentages for ion exchange, adsorption or reverse osmosis; the EPA fact sheet defers to its treatability database, which was not opened.
• Battery manufacturing and recycling effluent chemistry (lithium with fluoride, cobalt, nickel, sulfate) is not sourced here.
• Other GCC discharge standards were not read.
• No stoichiometry was read for the lithium selective sorbents (manganese and titanium oxide ion sieves, the lithium aluminate intercalation), so no equation is written for the adsorption row.

Sources

Identity

Name and symbol

Lithium, Li

Atomic number

3 protons

Position

group 1 · period 2 · s-block · alkali metal

CAS number

7439-93-2

Atomic structure

Atomic mass

6.941 u

Electron configuration

1s² 2s¹
[He] 2s¹

Electrons per shell

2, 1

Valence electrons

1 outer shell

Physical properties

State at room temperature

Solid

Melting point

453.65 K (180.5 °C)

Boiling point

1,615 K (1,341.85 °C)

Density

0.534 g/cm3

Appearance

silvery-white

Thermal conductivity

84.8 W/(m·K)

Electrical resistivity

92.8 nΩ·m at 20 °C

Electrical conductivity

10.78 MS/m

Crystal structure

body-centered cubic

Molar heat capacity

24.86 J/(mol·K)

Chemical properties

Oxidation states

+1

Electronegativity

0.98 (Pauling Scale)

Ionisation energy

5.392 eV
1st 520.2, 2nd 7,298.1, 3rd 11,815 kJ/mol

Electron affinity

0.618 eV

Atomic radius

empirical 128, covalent 128, van der Waals 182 pm

Ionic radius

Li⁺ 76 pm

Reactivity

The lightest alkali metal, with a single 2s electron that is easily lost; highly reactive and flammable, stored under oil or inert gas, yet the least vigorous member of its group.

with water

Reacts steadily with cold water, floating and fizzing, less violently than sodium, to lithium hydroxide and hydrogen: $\ce{2 Li + 2 H2O -> 2 LiOH + H2}$

with oxygen, air

Tarnishes quickly in moist air to a black crust of hydroxide, nitride and carbonate; burns in air to the oxide with some peroxide, and is the only alkali metal that also forms a nitride in air: $\ce{4 Li + O2 -> 2 Li2O}$

with acids

Reacts violently with dilute acids, faster than with water, giving the lithium salt and hydrogen.

with halogens

Reacts vigorously with all the halogens to salt-like lithium halides: $\ce{2 Li + Cl2 -> 2 LiCl}$

Typical compounds

• Li₂CO₃ lithium carbonate the key lithium compound; ore product, batteries, bipolar drug
• LiOH lithium hydroxide strong base; the black tarnish lithium forms in moist air
• LiCl lithium chloride very hygroscopic; the molten salt for lithium electrolysis
• Li₂O lithium oxide white oxide; special glasses and glass ceramics
• LiH lithium hydride ionic hydride used to store hydrogen
• LiAlH₄ lithium aluminium hydride reducing reagent in organic synthesis

Occurrence, production and use

Crustal abundance

2.0×101 milligrams per kilogram

Oceanic abundance

1.8×10-1 milligrams per liter

Occurrence and sources

It does not occur freely in nature; combined, it is found in small units in nearly all igneous rocks and in many mineral springs. Lepidolite, spodumene, petalite, and amblygonite are the more important minerals containing it.

Lithium is presently being recovered from brines of Searles Lake, in California, and from those in Nevada. Large deposits of quadramene are found in North Carolina. The metal is produced electrolytically from the fused chloride. Lithium is silvery in appearance, much like Na, K, and other members of the alkali metal series. It reacts with water, but not as vigorously as sodium. Lithium imparts a beautiful crimson color to a flame, but when the metal burns strongly, the flame is a dazzling white.

• dissolved in seawater about 0.18 mg/L; crustal estimate 20 mg/kg (Jefferson Lab figures via PubChem)
• spodumene, petalite, lepidolite, amblygonite (lithium aluminosilicates and phosphates) pegmatites and other igneous rocks; nine mineral operations in Australia and others in Brazil, Canada, China, Zimbabwe, Namibia and Portugal accounted for most hard-rock production in 2024
• lithium in continental brines salt flats of Argentina and Chile (four and two brine operations in 2024), five brine operations in China, geothermal and oilfield brines; Searles Lake and Nevada brines in the United States
• lithium clays (hectorite, claystone) Mexico and the United States, in development

Extraction, production

Brine evaporation and precipitation of lithium carbonate with sodium carbonate

$\ce{2 LiCl (aq) + Na2CO3 (aq) -> Li2CO3 (s) + 2 NaCl (aq)}$

The RSC states that brines yield lithium carbonate when treated with sodium carbonate; no equation is printed. Downstream, US producers convert domestic or imported carbonate, chloride and hydroxide into a range of lithium compounds (USGS).

Electrolysis of molten lithium chloride with potassium chloride to lithium metal

$\ce{2 LiCl (l) -> 2 Li (l) + Cl2 (g)}$

The metal is produced electrolytically from the fused chloride, first done in bulk by Bunsen and Matthiessen in 1855; the sources give the route in words only.

Uses

Many uses have been found for lithium and its compounds. Lithium has the highest specific heat of any solid element and is used in heat transfer applications. It is used to make special glasses and ceramics, including the Mount Palomar telescope's 200 inch mirror. Lithium is the lightest known metal and can be alloyed with aluminium, copper, manganese, and cadmium to make strong, lightweight metals for aircraft. Lithium hydroxide (LiOH) is used to remove carbon dioxide from the atmosphere of spacecraft. Lithium stearate (LiC18H35O2) is used as a general purpose and high temperature lubricant. Lithium carbonate (Li2CO3) is used as a drug to treat manic depression disorder.

Lithium reacts with water, but not as violently as sodium.

Since World War II, the production of lithium metal and its compounds has increased greatly. Because the metal has the highest specific heat of any solid element, it has found use in heat transfer applications; however, it is corrosive and requires special handling. The metal has been used as an alloying agent, is of interest in synthesis of organic compounds, and has nuclear applications. It ranks as a leading contender as a battery anode material as it has a high electrochemical potential. Lithium is used in special glasses and ceramics. The glass for the 200-inch telescope at Mt. Palomar contains lithium as a minor ingredient. Lithium chloride is one of the most hygroscopic materials known, and it, as well as lithium bromide, is used in air conditioning and industrial drying systems. Lithium stearate is used as an all-purpose and high-temperature lubricant. Other lithium compounds are used in dry cells and storage batteries. Lithium carbonate is used for the treatment of bipolar disease and other mental illness conditions.

• Batteries: rechargeable lithium-ion batteries for electric vehicles, portable electronics, tools and grid storage; primary batteries for pacemakers, toys and clocks batteries 87 percent of global lithium end use in 2024; global consumption about 220,000 tonnes, up 29 percent on 2023
• Glass and ceramics: lithium oxide in special glasses and glass ceramics; mineral concentrates used directly as flux ceramics and glass 5 percent of global end use in 2024
• Chemicals: lithium stearate as an all-purpose and high-temperature grease; lithium chloride and bromide as desiccants in air conditioning and industrial drying; continuous casting mould flux powders; aluminium-lithium and magnesium-lithium alloys for aircraft and armour lubricating greases 2 percent, air treatment 1 percent, casting mould flux 1 percent of global end use in 2024
• Pharmaceuticals: lithium carbonate as a drug for bipolar disorder medical 1 percent of global end use in 2024

Safety, toxicity

GHS classification, signal word Danger

• H260 In contact with water releases flammable gases which may ignite spontaneously Substances and mixtures which in contact with water, emit flammable gases
• H314 Causes severe skin burns and eye damage Skin corrosion/irritation
• H318 Causes serious eye damage Serious eye damage/eye irritation
• H371 May cause damage to organs Specific target organ toxicity, single exposure

Discovery and name

Discovered by

Johan August Arfwedson

Discovered

1817

First isolated

William Thomas Brande

Named by

not in sources

Origin of the name

from the Greek word λιθoς, stone