Lithium

Lithium is the chemical element with atomic number 3 and the symbol Li. It is a soft, silvery alkali metal in group 1 of the periodic table, and it has the lowest density of any metal. Lithium is best known today as the key ingredient in rechargeable batteries, although it also has long-standing uses in glass, ceramics, lubricants and medicine.

History of lithium

The first lithium mineral, petalite, was found on the Swedish island of Utö in the 1790s by the Brazilian naturalist Jozé Bonifácio de Andrada e Silva. He noticed that it burned with an intense crimson flame when thrown on a fire. In 1817 the Swedish chemist Johan August Arfvedson analyzed the mineral in Stockholm and concluded that it contained a previously unknown metal, which he recognized as a lighter relative of sodium.

The name comes from the Greek word lithos, meaning stone. Lithium was thought at the time to exist only in minerals, whereas the other alkali metals had first been found in plant material. Arfvedson could not separate the metal itself. William Brande obtained a tiny amount by electrolysis in 1821, too little to measure, and the German chemist Robert Bunsen and the British chemist Augustus Matthiessen finally produced it in bulk in 1855 by passing an electric current through molten lithium chloride.

Properties of lithium

Lithium has a standard atomic weight of about 6.94, and its single outer electron sits in the second shell, giving the electron configuration [He] 2s1. Under normal conditions it is a solid with a density of 0.534 grams per cubic centimeter, roughly half that of water, so a piece of lithium floats. It melts at 180.5 degrees Celsius and boils at 1,342 degrees Celsius.

Like the other alkali metals, lithium reacts with water, but less violently than sodium does. It also gives a flame a bright crimson color, which chemists use as a simple test. Burning lithium metal in air produces a dazzling white flame. Because the metal tarnishes and reacts readily, it is stored away from air and moisture and does not occur uncombined in nature.

PubChem lists a first ionization energy of 5.392 electronvolts for lithium and an electronegativity of 0.98 on the Pauling scale. Both numbers are low, which means the outer electron is held loosely and the atom easily forms a positive ion with a charge of plus one. This is the basis of its chemistry, including its behavior inside a battery.

Occurrence and production of lithium

Lithium makes up only a small fraction of the Earth’s crust, on the order of 20 milligrams per kilogram, but it is found in small amounts in nearly all igneous rocks and in the waters of many mineral springs. The important lithium-bearing minerals are spodumene, petalite, lepidolite and amblygonite. Brines, which are salty underground waters rich in dissolved minerals, are a second major source.

Most lithium is currently produced in Chile from brines that yield lithium carbonate when treated with sodium carbonate. Other recovery has taken place from brines in Nevada and from Searles Lake in California. To make the metal itself, producers pass electricity through a molten mixture of lithium chloride and potassium chloride, a process called electrolysis.

Uses of lithium

The most important use of lithium is in rechargeable batteries for mobile phones, laptops, digital cameras and electric vehicles. Lithium is also found in some non-rechargeable batteries, such as those in heart pacemakers, toys and clocks. The metal suits batteries because it has a high electrochemical potential, meaning it readily gives up electrons and so can store a lot of energy for its weight.

Lithium alloys with aluminum and magnesium to make strong, light materials. Aluminum-lithium alloys appear in aircraft, bicycle frames and high-speed trains, while a magnesium-lithium alloy has been used for armor plating. Lithium oxide goes into special glasses and glass ceramics, and the glass for the 200-inch mirror of the Palomar telescope contains lithium as a minor ingredient.

Several compounds have specialized roles. Lithium chloride and lithium bromide are strongly hygroscopic, meaning they absorb water vapor, so they are used in air conditioning and industrial drying systems. Lithium stearate is an all-purpose, high-temperature lubricant. Lithium hydroxide removes carbon dioxide from the air in spacecraft, and lithium hydride has been studied as a way to store hydrogen. Lithium carbonate is also used as a medicine for bipolar disorder, a condition formerly called manic depression, although the way it acts on the brain is still not fully understood.

The metal also has the highest specific heat of any solid element, so it can absorb a great deal of heat for its mass. That has made it of interest for heat transfer systems, although molten lithium is corrosive and needs special handling. Since World War II the production of lithium metal and its compounds has grown greatly.

Isotopes of lithium

Lithium has two stable isotopes, lithium-6 and lithium-7. According to IUPAC data, lithium-7 makes up about 92.41 percent of natural lithium and lithium-6 about 7.59 percent. Natural samples vary somewhat in this ratio, and scientists use the differences to trace the source of dissolved lithium in water, for example to tell water from marine sedimentary rocks from water altered by hot igneous processes.

Each isotope has its own industrial role. Lithium-7, as lithium hydroxide, helps control the acidity of the coolant in pressurized water nuclear reactors. Lithium-6 can capture a neutron to form helium and tritium, which makes it a fusion fuel material. During the 1950s large amounts of lithium-6 were removed from laboratory reagents for weapons work, so some commercial lithium sold afterward was depleted in that isotope.

Health and environmental effects of lithium

Lithium has no known biological role in humans, and the Royal Society of Chemistry describes it as toxic except in very small doses. Lithium-containing medicines are prescribed and monitored by physicians, and nothing here is guidance for personal use. Because lithium compounds are mined from brines and rock, extraction raises questions about water use and land disturbance in the regions where it is produced.

Editorially reviewed by Specialty Digest Editorial TeamLast reviewed November 4, 2026Researched and drafted with AI assistanceReport an issue