Lithium

Lithium
Atomic number	3
Symbol	Li
Group	1 (alkali metals)
Boiling point	1342 °C
Electronegativity	0.98 (Pauling)
Electron configuration	[He] 2s1
Melting point	180.5 °C
Period	2
Main isotopes	6Li, 7Li
Oxidation states	+1
Phase STP	Solid
Wikidata	Q568

Lithium (Li, atomic number 3) is a silvery alkali metal in group 1 of the periodic table, with atomic weight about 6.94 It is the lightest of all solid elements and the least dense metal; it floats on water and tarnishes in air In its pure form lithium is a soft, silvery-white metal that readily loses its single valence electron (configuration [He] 2s¹) to form the Li⁺ cation These basic facts – symbol Li, Group 1, period 2, s-block – underlie the element’s characteristic chemistry. A summary of key numerical properties (atomic number, weight, melting/boiling points, density, electron configuration, etc.) is given in the data table below. In the rest of this article we will survey lithium’s atomic structure, isotopes, phases and compounds, physical behavior, chemical reactivity, natural occurrence and production, technological applications, biological effects and safety, and its history and name.

Atomic Structure and Electron Configuration

Lithium atoms contain 3 protons and 3 electrons, with the ground-state electronic configuration [He] 2s¹. In other words, two electrons fill the 1s core and one occupies the 2s valence orbital This half-filled 2s shell means Li has only one valence electron (in contrast with H, which has a 1s electron). As an alkali metal, Li’s single valence electron is relatively loosely held, but compared to other alkali metals Li’s first ionization energy is higher and its electronegativity is larger due to its small size. On the Pauling scale lithium’s electronegativity is about 0.98 higher than for sodium or potassium but still quite low overall. The first ionization energy is roughly 520 kJ/mol (≈5.39 eV), reflecting that the small Li atom holds its sole electron more tightly than do the heavier alkalis. Conversely, lithium atoms are very small: the covalent radius is only on the order of 150 pm. Small size and a single valence electron give lithium high metallic character and strong electropositivity, but also slightly higher ionization energy and electronegativity than its heavier group-1 neighbors. In periodic trends, lithium as the first alkali metal is an outlier closer to beryllium; it does not quite behave identically to Na or K in all ways (the so-called diagonal relationship with Mg is one example).

Isotopes and Nuclear Properties

Naturally occurring lithium consists almost entirely of two stable isotopes: Li-6 and Li-7. Li-7 is the more abundant, composing about 92.5% of natural lithium, while Li-6 makes up the remaining 7.5% or so (This isotopic mix yields the standard atomic weight ≈6.94.) Both Li-6 and Li-7 have anomalously low binding energies for their nuclei, a legacy of Big Bang nucleosynthesis. Indeed, lithium (along with hydrogen and helium) is one of the few primordial elements formed in stellar Big Bang processes Heavier radioactive lithium isotopes (Li-8, Li-9, Li-11, etc.) are produced in nuclear reactions but decay rapidly (Li-8 has half-life <1 s, Li-11 ~9 ms) and are not found naturally.

Li-6 and Li-7 also differ in nuclear spin: Li-6 has spin 1, while Li-7 has spin 3/2. These spins enable useful NMR (nuclear magnetic resonance) studies of lithium compounds, especially Li-7 NMR in battery materials. In nuclear technology, Li-7 (which has very low neutron absorption cross-section) is important for reactor chemistry, whereas Li-6 is valuable for tritium breeding. When Li-6 absorbs a neutron it can undergo nuclear reaction producing tritium (^3H) and helium In practice, enriched Li-6 is used in reactors or fusion devices to generate tritium fuel. For example, in hydrogen (thermonuclear) bombs lithium deuteride (Li-6D) is fused to produce copious tritium, fueling explosive fusion reactions Li-7 is also added as LiOH to pressurized water reactor (PWR) coolant and used in molten-salt reactor coolants, because it helps stabilize pH while remaining essentially transparent to neutrons In short, lithium’s nuclides have unique nuclear roles: Li-7 for corrosion control in reactors, and Li-6 for generating fusion fuel.

Allotropes and Typical Compounds

Lithium metal has no allotropes in the usual sense – there is just the metallic form. At room conditions lithium crystallizes in a body-centered cubic (bcc) structure (Under high pressure it transforms to a close-packed phase – researchers report a bcc→fcc transition around ~7 GPa Unlike carbon or phosphorus, lithium does not form distinct non-metallic forms.

However, lithium forms a wide variety of compounds. In nearly all of them lithium is in oxidation state +1 (Li⁺), consistent with losing that one valence electron. Some important inorganic classes include:

Oxides/hydroxides: Lithium oxide (Li₂O) is the normal alkali-metal oxide; lithium peroxide (Li₂O₂) also occurs, as does superoxide (LiO₂) in special contexts. Lithium hydroxide (LiOH) is a strong base used industrially (e.g. in batteries or CO₂ scrubbing).
Nitrides: Lithium nitride Li₃N is notable as the only simple alkali nitride. Solid Li reacts with N₂ to form this red nitride compound
Halides: All lithium halides LiF, LiCl, LiBr, LiI are stable white solids. Lithium chloride and bromide are extremely hygroscopic (they absorb moisture) and are used in air-conditioning dehumidifiers. Lithium fluoride is less soluble and is used in special glasses and fluxes.
Cations and salts: Lithium forms soluble salts like Li₂CO₃ (carbonate), LiNO₃, Li₂SO₄, etc. Lithium carbonate is a major commercial chemical (extracted from brines) and also a pharmaceutical (see below). Sulfides such as Li₂S are used in some batteries.
Metal hydride: Lithium hydride (LiH) accepts and stores hydrogen – it is used as a hydrogen-storage and fusion-fuel material. Indeed, LiH (with Li-6) is the fuel in many thermonuclear devices
Organolithium: Lithium uniquely forms organometallic compounds in which Li is covalently bonded to carbon. These organolithium reagents (for example n-butyllithium, C₄H₉Li) are extremely useful strong bases/nucleophiles in synthetic chemistry Butyllithium (C₄H₉Li), for instance, is made by reacting butyl bromide with lithium metal and is widely used in industry (e.g. in making synthetic rubber)

Lithium compounds often display ionic bonding with Li⁺, but being the smallest alkali cation, Li⁺ can also form well-defined coordination complexes with oxygen- or nitrogen-donor ligands (as in cryptates or crown-ether complexes). In many cases the properties of Li compounds (solubility, reactivity) differ markedly from their sodium or potassium analogs because of Li’s small size and high polarization power.

Physical Properties

At standard conditions lithium is a solid metal with density about 0.534 g/cm³ (the lowest of all the metals). Its melting point is relatively low (180.5 °C) and boiling point high (about 1342 °C) Structurally, solid lithium is body-centered cubic with one atom per corner of the cube plus one in the center Lithium is soft: it can be cut with a knife (Mohs hardness ~0.6) and malleable, though slightly harder than sodium. The metal’s appearance is lustrous silver when freshly cut, but it quickly tarnishes to a dull gray in air due to oxide formation.

Lithium is an excellent conductor of heat and electricity (as typical for a metal), though less so than copper or aluminum on a per-weight basis (because it has very low density). Its specific heat (~3.58 J/g·K) is exceptionally high (higher than any other metal), meaning it absorbs a lot of heat before warming. Unusually for a metal, lithium has been observed to become superconducting under certain conditions at very low temperatures (around 0.5 K).

Spectroscopically, lithium exhibits a distinctive crimson-red emission line when excited. In a flame test or emission spectrum the strongest line is around 670.8 nm (red). Indeed, “lithium salts impart a crimson color to a flame,” a fact used for identification (By contrast, sodium gives a yellow flame, potassium violet, etc.) The strong red flame is a classic qualitative test for lithium.

Thermodynamically, lithium metal has a very negative standard reduction potential (Li⁺ + e⁻ → Li⁰ occurs at –3.040 V versus SHE), reflecting its tendency to oxidize and form Li⁺. In molten or aqueous conditions Li⁺ is relatively small (ionic radius ~90 pm for six-fold coordination) and highly polarizing. Among its physical magnetism, metallic lithium is paramagnetic in bulk. In many properties (thermal expansion, compressibility, speed of sound) lithium fits general metal trends for a light alkali.

Chemical Reactivity and Trends

Like other alkali metals, lithium is highly chemically reactive. It readily loses its single valence electron to form Li⁺ and undergo redox reactions. A famous reaction is with water: lithium floats on water and reacts vigorously (though somewhat less violently than sodium or potassium) to yield lithium hydroxide and hydrogen gas

2 Li (s) + 2 H₂O (l) → 2 LiOH (aq) + H₂ (g).

The dilute LiOH solution produced is strongly basic. This vigorous reaction is so characteristic that lithium metal is usually stored under mineral oil or kerosene to keep it from moisture Unlike sodium, lithium does not form a stable Li⁻ (lithide) anion in solution; it exclusively forms Li⁺.

Lithium also reacts with oxygen: it burns in air to form oxides (Li₂O, Li₂O₂). It does not react with nitrogen to the extent one might expect, except at high temperature or with sparks: lithium metal exposed to N₂ gas will gradually form lithium nitride Li₃N a red solid. This nitride formation is unusual since other alkali metals do not form M₃N. At moderately high temperatures, Li will react with halogens to make LiX (e.g. LiCl, LiBr), and with hydrogen to make lithium hydride (LiH).

Many lithium compounds are stable and colorless (halides, most salts). Lithium oxide (Li₂O) and hydroxide are strongly basic. Lithium carbonate (Li₂CO₃) and lithium hydroxide are similar to beryllium hydroxide in chemistry (due to diagonal relationship), but like all alkalis they are water-soluble bases. Lithium metal itself is a strong reducing agent (e.g. it reduces metal halides in molten salts). For example, molten Li₂O can dissolve metals like copper as Cu and leave behind Li or produce Li₂O glasses. Lithium metal is not stored under water or alcohol because of violent reaction; it is usually kept under inert oil. Overall, lithium’s chemical behavior follows alkali trends (very electropositive, forming water-soluble +1 compounds), but with a few distinctive twists (its ability to form Li₃N, the importance of crescent-like small polarizing cation, etc.)

Occurrence and Production

Lithium is relatively rare in the universe because most is consumed in stars or was only sparsely produced in the Big Bang In the cosmos its abundance is much lower than that of other light elements like hydrogen, helium or carbon. In the Solar System, cosmic lithium is still trace: it is about 20–25 parts per billion by mass in the Sun and orders of magnitude less abundant than carbon or oxygen.

On Earth, lithium is more common (though still a trace element). Its abundance in the Earth’s crust is on the order of tens of parts per million: it is present in nearly all igneous rocks and certain clays and soils Major minerals include spodumene (LiAlSi₂O₆), lepidolite, petalite (LiAl(SiO₄)O), and amblygonite. Lithium also occurs dissolved in ocean water (around 0.18 ppm in seawater) and in mineral brines (saltwater deposits). Bolivia’s Uyuni salt flat, Chile’s Atacama salt lake, and Argentina’s salars contain huge reserves of lithium-rich brine.

Commercial lithium is obtained either from evaporating brines or mining hard rock. Brine extraction (adept in South America) involves pumping lithium-bearing saltwater into shallow ponds and evaporating to precipitate Li₂CO₃ or LiCl. Hard-rock mining (Australia, China, U.S.) targets spodumene and other ores, which are crushed and chemically processed (often by treatment with acid or soda ash) to yield lithium chemicals. According to current data, about 60–70% of global lithium production comes from brines, the rest from ore. The major producing countries (as of the 2020s) are Australia (spodumene mining, ~50% of reserves), Chile and Argentina (lithium brines, ~30% combined), and China (~17% via both methods) The U.S. is a smaller producer. Bolivia has very high lithium reserves but has produced only a tiny fraction (pending development).

Worldwide lithium production has surged in recent years due to demand for batteries. The intermediate product of most processes is lithium carbonate (Li₂CO₃). Lithium carbonate is also converted to lithium hydroxide (LiOH) by treatment with lime, which is often the form needed for batteries. Metallic lithium is made by molten-salt electrolysis: typically LiCl (often mixed with KCl to lower the melting point) is electrolyzed to deposit pure Li at the cathode (a method analogous to producing Na metal). Laboratory-scale methods can produce small Li samples (e.g. Robert Bunsen and Matthiessen achieved bulk Li by electrolysis of molten LiCl in the 1850s but today’s industrial plants handle many tons.

Applications and Technology

By far the largest use of lithium today is in energy storage: rechargeable lithium-ion (Li-ion) batteries Li-ion batteries, which shuttle Li⁺ ions between anode and cathode materials, power most smartphones, laptops, digital cameras and increasingly electric vehicles and grid storage. Their high voltage (≈3–4 V per cell), long life and light weight have made Li-ion batteries ubiquitous. (Smaller disposable “Li-metal” batteries also exist for specialty uses, but these use lithium metal as an electrode, not as a moving ion.) The growth of Li-ion technology is a primary driver of lithium demand; collectively, batteries now consume over half of annual lithium supply

Lithium alloys and compounds have many other roles. Aluminum-lithium and magnesium-lithium alloys combine low weight with high strength and are used in aerospace, transportation and armor plating Lithium oxide and lithium fluoride (with other oxides/fluorides) are additives in specialty glasses and ceramics, improving heat resistance or refractive index. Lithium grease (thickened with lithium stearate) is a common automotive and industrial lubricant, valued for thermal stability Lithium fluxes (such as Li₂O in steelmaking) aid in metal refining.

In electronics, lithium compounds are key. For example, lithium tantalate (LiTaO₃) and niobate (LiNbO₃) are important piezoelectric materials. Lithium fluoride is used in HF optics. In nuclear technology, enriched Li-7 in the coolant (as LiOH or LiF) helps control reactor chemistry Lithium deuteride (rich in Li-6) has been a fusion fuel.

Lithium also has a long history in organic chemistry. As noted, organolithium reagents like n-butyllithium are fundamental in industrial synthesis, serving as very strong bases or nucleophiles In addition, lithium aluminium hydride (LiAlH₄) is a major reducing agent used to turn esters and carboxylic acids into alcohols These applications leverage the high reactivity of Li–C or Li–H bonds.

In summary, lithium’s technological roles are broad: batteries (the dominant use by far lightweight alloys, specialty ceramics/glasses, lubricants, and as a reagent in chemical synthesis. According to one review, more than three-quarters of lithium production today goes into batteries

Biology, Environment, and Safety

Lithium is not known to be an essential nutrient for any life form Nevertheless, trace amounts of lithium are present in soils, plants and animals because it is so ubiquitous. Higher animals and humans tolerate lithium; in fact, assorted foods (grains, vegetables, dairy) contain small lithium. There is no recognized biological role for Li₊, and it is not part of any known enzyme system.

Medically, lithium ions are used as a drug, famously as lithium carbonate (Li₂CO₃) to treat bipolar disorder (manic depression). Lithium salts stabilize mood in patients prone to mania and are one of the oldest psychiatric medications. In small controlled doses (hundreds of mg/day), Li₂CO₃ can be therapeutic However, lithium has a narrow therapeutic index: blood levels must be monitored, because too much lithium causes toxicity (nausea, tremors, kidney damage, etc.). Indeed, toxic poisoning cases usually involve LiCl or Li₂CO₃ overdoses. But at normal therapeutic or environmental levels, lithium is generally well tolerated.

Environmental cycling of lithium is complex. Because lithium salts are soluble, lithium tends to remain in water and soil rather than bioaccumulate. Coal-burning power plants, for instance, release trace lithium into the air. Lithium concentrations can become elevated near mining or improper disposal sites; batteries and brines discharges are potential sources of pollution. In nature, lithium is transported by rivers to oceans (sea water ~0.17 ppm Li) and eventually deposited in marine sediments. Lithium has a relatively long residence time (~1 million years) in the oceans. Overall, lithium is considered of low ecotoxicological concern compared to heavy metals, but high concentrations (from spills or waste) can harm aquatic life.

Safety-wise, lithium metal and compounds require careful handling. Lithium metal is highly reactive and flammable: it ignites in air if unprotected and reacts violently with water (releasing flammable H₂) Salt or earth fire extinguishers are needed if lithium metal burns (water or CO₂ extinguishers can make it worse). Consequently, elemental Li is stored under oil or inert atmosphere. Many lithium compounds (like LiOH, Li₂O₂) are caustic bases; dust or solutions can burn skin or eyes. Workers are advised to avoid inhaling lithium dust or fumes: OSHA regulates Li compound exposures at low ppm levels. Lithium salts in batteries pose a fire hazard if cells short or overheat, which is a modern safety challenge in electronics.

Bio-safety: Lithium ions are toxic to aquatic organisms at high levels, but the general threshold is much higher than therapeutic plasma levels in humans. In humans, occupational exposure limits (for Li salts) are around 0.1–0.8 mg/m³ (as dust), so normal lab or mining exposure is carefully controlled. Environmental guidelines do not typically include "safe drinking level" for lithium (it is not regulated as a pollutant in drinking water), although some studies have looked at lithium’s subtle effects on ecosystems.

In sum, lithium must be handled with respect for its reactivity and toxicity, but it is not nearly as acutely poisonous as heavy metals. For example, its use in medicine and industry (including in animals via medication) demonstrates that reasonable exposures can be managed safely. Still, direct contact with Li metal or concentrated Li compounds is hazardous. As with any corrosive or pyrophoric metal, proper protective equipment (gloves, goggles) and protocols are essential.

History and Etymology

Lithium’s story begins in the late 18th and early 19th centuries with European mineralogy. In 1800 a Swedish chemist, Jöns Jacob Berzelius, discovered potassium in an unusual mineral from the Swedish island of Utö. Building on that, in 1817 (by some accounts 1807) Johan August Arfwedson, working in Stockholm, analyzed the petalite mineral (LiAl(Si₂O₄)O) and found an unknown alkali. He concluded it contained a new element, which he called “lithion” after the Greek lithos (stone), contrasting with potash (from plants) which gave sodium and potassium The name “lithium” (Latinized) stuck.

However, Arfwedson did not isolate the pure metal. It was only in 1855 that Robert Bunsen and Augustus Matthiessen first prepared bulk lithium metal by electrolyzing molten lithium chloride Thus lithium became the last of the fundamental alkali metals to be isolated. Earlier, Sir Humphry Davy had isolated sodium and potassium in 1807 by electrolysis of hydroxides, but lithium required a lower-melting electrolyte. Bunsen and Matthiessen’s work completed the picture of the alkali group.

Later milestones include William Cruickshank’s identification (around 1818) of lithium’s characteristic red flame color, and the development of analytical techniques to measure lithium in minerals and brines. In 1923, the nuclear physicist Arthur E. Ruark gave neutron scattering evidence that lithium’s atomic mass had to be about 7, confirming the existence of two isotopes.

In the 20th century lithium found new roles. In 1948 Australian psychiatrist John Cade discovered lithium carbonate’s calming effect on manic patients, leading to its use in treating bipolar disorder. In the 1950s lithium compounds entered nuclear weapons design for thermonuclear bombs. The lithium-ion battery was first proposed by chemist M. Stanley Whittingham in 1976 and later commercialized by John Goodenough and colleagues; by the 1990s Sony and others were producing Li-ion cells for consumer electronics. In 2010s electric vehicles gave lithium a new surge of importance. Throughout, lithium’s name has remained “stone,” a reminder of its origin in the mineral world, even though today we mine vast brines and pegmatites to extract it.

The data table below summarizes lithium’s key identifiers and constants for quick reference (isotopic abundances, atomic data, physical constants, etc.).

Property	Value
Symbol	Li
Atomic number (Z)	3
Standard atomic weight	6.94 (6.941)
Electron configuration	[He] 2s¹
Group, Period, Block	1, 2, s-block (alkali metal)
State at STP	Solid
Density (20 °C)	0.534 g·cm⁻³
Melting point	180.50 °C
Boiling point	1342 °C
Crystal structure	Body-centered cubic (α-Li)
Electronegativity (Pauling)	0.98
First ionization energy	520 kJ·mol⁻¹ (5.39 eV)
Oxidation state	+1 (typical)
Common isotopes	⁶Li (7.59%), ⁷Li (92.41%)
CAS Registry Number	7439-93-2
Flame test color	Crimson red