Caesium chloride

The caesium chloride structure adopts a primitive cubic lattice with a two-atom basis, where both atoms have eightfold coordination. The chloride atoms lie upon the lattice points at the corners of the cube, while the caesium atoms lie in the holes in the center of the cubes; an alternative and exactly equivalent 'setting' has the caesium ions at the corners and the chloride ion in the center. This structure is shared with CsBr and CsI and many binary metallic alloys. In contrast, the other alkaline halides have the sodium chloride (rocksalt) structure. When both ions are similar in size (Cs+ ionic radius 174 pm for this coordination number, Cl− 181 pm) the CsCl structure can be adopted, when they are different (Na+ ionic radius 102 pm, Cl− 181 pm) the sodium chloride structure is adopted. Upon heating to above 445 °C, the normal caesium chloride structure (α-CsCl) converts to the β-CsCl form with the rocksalt structure (space group Fmm). ==Physical properties==

Caesium chloride is colorless in the form of large crystals and white when powdered. It readily dissolves in water with the maximum solubility increasing from 1865 g/L at 20 °C to 2705 g/L at 100 °C. The crystals are highly hygroscopic and deliquescent. Caesium chloride crystals gradually disintegrate at ambient conditions. sulfur dioxide (2.95 g/L at 25 °C), ammonia (3.8 g/L at 0 °C), acetone (0.004% at 18 °C), acetonitrile (0.083 g/L at 18 °C), ethylacetate and other complex ethers, butanone, acetophenone, pyridine and chlorobenzene. Despite its wide band gap of about 8.35 eV at 80 K, caesium chloride weakly conducts electricity, and the conductivity is not electronic but ionic. The conductivity has a value of the order 10−7 S/cm at 300 °C. It occurs through nearest-neighbor jumps of lattice vacancies, and the mobility is much higher for the Cl− than Cs+ vacancies. The conductivity increases with temperature up to about 450 °C, with an activation energy changing from 0.6 to 1.3 eV at about 260 °C. It then sharply drops by two orders of magnitude because of the phase transition from the α-CsCl to β-CsCl phase. The conductivity is also suppressed by application of pressure (about 10 times decrease at 0.4 GPa) which reduces the mobility of lattice vacancies. ==Reactions==

Caesium chloride completely dissociates upon dissolution in water, and the Cs+ cations are solvated in dilute solution. CsCl converts to caesium sulfate upon being heated in concentrated sulfuric acid or heated with caesium hydrogen sulfate at 550–700 °C: :2 CsCl + H2SO4 → Cs2SO4 + 2 HCl :CsCl + CsHSO4 → Cs2SO4 + HCl Caesium chloride forms a variety of double salts with other chlorides. Examples include 2CsCl·BaCl2, 2CsCl·CuCl2, CsCl·2CuCl and CsCl·LiCl, and with interhalogen compounds: : ==Occurrence and production==

s. Caesium chloride occurs naturally as an impurity in the halide minerals carnallite (KMgCl3·6H2O with up to 0.002% CsCl), : Only about 20 tonnes of caesium compounds, with a major contribution from CsCl, were being produced annually around the 1970s and is sold internationally through a UK dealer. The salt is synthesized at 200 °C because of its hygroscopic nature and sealed in a thimble-shaped steel container which is then enclosed into another steel casing. The sealing is required to protect the salt from moisture. Laboratory methods In the laboratory, CsCl can be obtained by treating caesium hydroxide, carbonate, caesium bicarbonate, or caesium sulfide with hydrochloric acid: :CsOH + HCl → CsCl + H2O :Cs2CO3 + 2 HCl → 2 CsCl + 2 H2O + CO2 ==Uses==

Precursor to Cs metal Caesium chloride is the main precursor to caesium metal by high-temperature reduction: Nuclear medicine and radiography Caesium chloride composed of radioisotopes such as 137CsCl and 131CsCl, and screens of cathode ray tubes. Therefore, it can be useful in electrophysiology experiments in neuroscience. ==Toxicity==

Caesium chloride has a low toxicity to humans and animals. However, caesium chloride powder can irritate the mucous membranes and cause asthma. ==See also==