Cesium is a very soft, silvery-gold alkali metal. It has the lowest ionization energy of all elements, melts just above room temperature, and reacts violently with water and air.
The ground-state configuration is [Xe] 6s1. That single, loosely held 6s electron (very low first ionization energy) is readily lost to form Cs+, making Cs the most reactive stable alkali metal under standard conditions.
Cs has relatively weak metallic bonding because its valence 6s electron is diffuse and poorly overlaps with neighbors. This lowers lattice cohesion, giving a melting point of about 28.5 °C, so warm rooms can partially liquefy Cs.
Caesium reacts violently with water, producing caustic cesium hydroxide and hydrogen gas; the heat can ignite the H2:
\(\mathrm{2\,Cs(s) + 2\,H_2O(l) \rightarrow 2\,CsOH(aq) + H_2(g)}\)
In air, Cs rapidly forms oxides/peroxides/superoxides (e.g., superoxide \(\mathrm{CsO_2}\)) and must be stored under dry, inert conditions.
Key examples include:
The SI second is defined by the hyperfine transition frequency of the ground state of \(^{133}\mathrm{Cs}\): 9,192,631,770 Hz. This microwave transition is extremely reproducible, enabling clocks with uncertainties better than 10−16 in advanced systems.
Caesium imparts a blue-violet (sometimes described as violet with blue tones) color to a flame due to characteristic electronic transitions of Cs+.
Store Cs under dry, oxygen-free mineral oil or inert gas (argon) in sealed ampoules. Use inert-atmosphere techniques, full face/eye protection, gloves, and non-sparking tools. Keep completely away from water, moist air, and oxidizers.
Cs is obtained from pollucite (a cesium aluminosilicate ore). Processing involves alkaline digestion and selective precipitation/ion exchange to isolate Cs salts, which can be reduced (e.g., with calcium) to the metal.
Cesium hydroxide is a strong base that fully dissociates in water and has very high basicity due to the large, weakly solvated Cs+ cation, leaving OH− highly active. It can rapidly saponify organics and damage glass at elevated temperatures.
Cs combines directly with halogens to give ionic halides:
\(\mathrm{2\,Cs(s) + Cl_2(g) \rightarrow 2\,CsCl(s)}\)
Here Cs is oxidized to Cs+ (from 0 to +1) and chlorine is reduced to Cl− (from 0 to −1).