Lithium (Li), a soft, silvery-white alkali metal, is the lightest metal and holds an important position in modern technology, particularly in energy storage. An isotope is defined by the number of neutrons in an atom’s nucleus; atoms of the same element have the same number of protons but varying numbers of neutrons. The total count of lithium isotopes depends on whether one considers only the stable forms found in nature or includes the highly unstable varieties synthesized in laboratories.
The Two Stable Forms of Lithium
Lithium found naturally on Earth primarily consists of two stable isotopes: Lithium-6 (\(^6\)Li) and Lithium-7 (\(^7\)Li). Both isotopes share the atomic number 3, meaning their nuclei each contain three protons. The difference lies in their neutron count, with \(^6\)Li having three neutrons and \(^7\)Li having four neutrons.
The vast majority of terrestrial lithium is the heavier isotope, \(^7\)Li, which accounts for approximately 92.4% of all naturally occurring lithium atoms. The lighter isotope, \(^6\)Li, makes up the remaining 7.6%. These percentages, known as natural abundances, determine the standard atomic weight of lithium listed on the periodic table.
Natural abundances can vary slightly depending on the source material, as geological and industrial processes can cause minor isotopic separation. For instance, certain commercial lithium supplies may show a depletion of \(^6\)Li due to historical separation for nuclear applications. Despite these minor variations, \(^6\)Li and \(^7\)Li are the only stable forms that persist in the environment.
The Known Unstable Varieties
Scientists have created several unstable isotopes of lithium in controlled laboratory settings. The known isotopes span from Lithium-4 (\(^4\)Li) to Lithium-12 (\(^{12}\)Li) and do not exist in nature for any significant duration. These isotopes have extremely short half-lives, meaning they decay almost instantly after formation.
The most stable radioactive isotope, Lithium-8 (\(^8\)Li), has a half-life of only about 838 milliseconds. Other forms, such as Lithium-11 (\(^{11}\)Li), known for its exotic “halo” nucleus, decay in mere milliseconds. The shortest-lived isotope, \(^4\)Li, decays in a fraction of a second, on the order of \(10^{-23}\) seconds. These fleeting isotopes are studied in nuclear physics to understand fundamental nuclear structure and forces.
Industrial Uses of Lithium Isotopes
While most commercial uses of lithium, such as in lithium-ion batteries, use the element in its natural isotopic mix, certain advanced technologies demand the isotopes be separated for specific roles. The distinct nuclear properties of \(^6\)Li and \(^7\)Li make them indispensable in the nuclear industry. This separation process, which is often difficult and costly, yields materials with purities exceeding 99% for each isotope.
Lithium-7 is highly valued for use in pressurized water nuclear reactors (PWRs). Lithium-7 hydroxide is added to the primary coolant water as a pH stabilizer to minimize corrosion in the reactor circuit. It is chosen because its nucleus has a very low tendency to absorb neutrons, an important feature in a reactor core. Lithium-7 is also a key component in the molten fluoride salts used in advanced molten salt reactors (MSRs).
Conversely, Lithium-6 is prized for its high affinity for absorbing neutrons. This property is crucial in the production of tritium, a radioactive isotope of hydrogen used as fuel in fusion energy research and in thermonuclear weapons. When \(^6\)Li is bombarded with neutrons, it undergoes a reaction that produces tritium and helium. Therefore, a separate supply of \(^6\)Li is necessary for developing future fusion power reactors.