Lithium

6.94

A lithium battery powers an electronic device.

Classification:	Lithium is an alkali metal
Atomic weight:	6.941 g/mol
Density @ 20 ^oC:	0.53 g/cm³
Atomic volume:	13.10 cm³/mol

Lithium was made in the first few minutes of the Universe, when temperatures everywhere were high enough for nuclear fusion to occur. This short, high energy phase is represented at the very bottom of the diagram. Image: Gnixon

Discovery of Lithium

Author: Dr. Doug Stewart

Lithium was discovered by Johan Arfvedson in 1817, during an analysis of petalite (LiAlSi₄O₁₀).

He found the petalite contained "silica, alumina and an alkali." ⁽¹⁾

The new alkali metal in the petalite had unique properties. It required more acid to neutralize it than sodium and its carbonate was only sparingly soluble in water - unlike sodium carbonate. The new alkali differed from potassium because it did not give a precipitate with tartaric acid.,

Arfvedson tried to produce a pure sample of the new metal by electrolysis, but he was unsuccessful; the battery he used was not powerful enough. ⁽²⁾

The pure metal was isolated the following year by both William Brande and Humphry Davy working independently.

Davy obtained a small quantity of lithium metal by electrolysis of lithium carbonate. ⁽³⁾

He noted the new element had a red flame color somewhat like strontium and produced an alkali solution when dissolved in water.

In days less safety-conscious than the present, Brande said of lithium, "its solution tastes acrid like the other fixed alkalies." ⁽⁴⁾

By 1855 Robert Bunsen and Augustus Matthiessen were independently producing the metal in large quantities by electrolysis of molten lithium chloride.

Lithium's name is derived from the Greek word 'lithos,' meaning, 'stone.'

Interesting Facts: Did You Know?

Interesting Facts about Lithium: Did You Know?

Lithium is believed to be one of only three elements - the others are hydrogen and helium - produced in significant quantities by the Big Bang.
Lithium is the only alkali metal that reacts with nitrogen.
Humphrey Davy produced some of the world's first lithium metal from lithium carbonate. Today lithium carbonate - or more precisely the lithium ions in lithium carbonate - are used to inhibit the manic phase of bipolar (manic-depressive) disorder.
Batteries based on lithium have revolutionized consumer devices such as computers and cell phones. For a given battery weight, lithium batteries pack a lot of energy compared with batteries based on other metals; in other words, lithium batteries have high energy density.

States

State (s, l, g): solid
Melting point: 453.69 K (180.54 ^oC)

Boiling point: 1615 K (1347 ^oC)

Energies

Specific heat capacity: 3.58 J g^-1 K^-1
Heat of fusion: 3.00 kJ mol^-1
1^st ionization energy: 520.2 kJ mol^-1
3^rd ionization energy: 11815.0 kJ mol^-1

Heat of atomization: 159 kJ mol^-1
Heat of vaporization: 147.1 kJ mol^-1
2^nd ionization energy: 7298.1 kJ mol^-1
Electron affinity: 59.63 kJ mol^-1

Oxidation & Electrons

Shells: 2,1
Minimum oxidation number: -1
Min. common oxidation no.: 0
Electronegativity (Pauling Scale): 0.98

Electron configuration: 1s² 2s¹
Maximum oxidation number: 1
Max. common oxidation no.: 1
Polarizability volume: 24.3 Å³

Appearance & Characteristics

Structure: bcc: body-centered cubic
Hardness: 0.6 mohs

How do you extract a strip of lithium from an energizer battery?

Powdered lithium reacts with water.

Color: silvery

Harmful effects:
Lithium is corrosive, causing skin burns as a result of the caustic hydroxide produced in contact with moisture. Women taking lithium carbonate for bi-polar disorder may be advised to vary their treatment during pregnancy as lithium may cause birth defects.

Characteristics:
Lithium is soft and silvery white and it is the least dense of the metals. It is highly reactive and does not occur freely in nature.

Freshly cut surfaces oxidize rapidly in air to form a black oxide coating. It is the only common metal (but see radium) that reacts with nitrogen at room temperature, forming lithium nitride.

Lithium burns with a crimson flame, but when the metal burns sufficiently well, the flame becomes a brilliant white.

Lithium has a high specific heat capacity and it exists as a liquid over a wide temperature range.

Uses:
Pure lithium metal is used in rechargeable lithium ion batteries and the metal is used as an alloy with aluminum, copper, manganese, and cadmium to make high performance aircraft parts.

Lithium also has various nuclear applications, for example as a coolant in nuclear breeder reactors and a source of tritium, which is formed by bombarding lithium with neutrons.

Lithium carbonate is used as a mood-stabilizing drug.

Lithium chloride and bromide are used as desiccants.

Lithium stearate is used as an all-purpose and high-temperature lubricant.

Reactions & Compounds

Reaction with air: vigorous,⇒ Li₂O
Reaction with 15 M HNO₃: vigorous,⇒ LiNO₃
Oxide(s): Li₂O
Hydride(s): LiH

Reaction with 6 M HCl: vigorous,⇒ H₂, LiCl
Reaction with 6 M NaOH: mild, ⇒ H₂, LiOH
Chloride(s): LiCl

Radius

Atomic radius: 145 pm
Ionic radius (2+ ion): pm
Ionic radius (2- ion): pm

Ionic radius (1+ ion): 90 pm
Ionic radius (3+ ion): pm
Ionic radius (1- ion): pm

Conductivity

Thermal conductivity: 84.8 W m^-1 K^-1

Electrical conductivity: 11.7 x 10⁶ S cm^-1

Abundance & Isotopes

Abundance earth's crust: 20 parts per million by weight, 60 parts per million by moles
Abundance solar system: 60 parts per trillion by weight, 10 parts per trillion by moles
Cost, pure: $27 per 100g
Cost, bulk: $9.50 per 100g

Source: Lithium does not occur as a free element in nature. It is found in small amounts in ores from igneous rocks and in salts from mineral springs. Pure lithium metal is produced by electrolysis from a mixture of fused (molten) lithium chloride and potassium chloride.

Isotopes: Lithium has 7 isotopes whose half-lives are known, with mass numbers 5 to 11. Of these, two are stable: ⁶Li and ⁷Li.

References

1. Thomas Thomson, A system of chemistry of inorganic bodies, 1831, Volume 1
2. Mary Elvira Weeks, Discovery of the Elements., 2003, p125 Kessinger Publishing.
3. Royal Institution of Great Britain, The Quarterly Journal of Science and the Arts, Volume 5, 1818 p338. (pdf)
4. William Thomas Brande, William James MacNeven A manual of chemistry (1821) p191. (pdf)

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