Caesium chloride or cesium chloride is the inorganic compound with the formula CsCl. This colorless salt is an important source of caesium ions in a variety of niche applications. Its crystal structure forms a major structural type where each caesium ion is coordinated by 8 chloride ions. Caesium chloride dissolves in water. CsCl changes to NaCl structure on heating. Caesium chloride occurs naturally as impurities in carnallite (up to 0.002%), sylvite and kainite. Less than 20 tonnes of CsCl is produced annually worldwide, mostly from a caesium-bearing mineral pollucite.^[7]

Caesium chloride is widely used in isopycnic centrifugation for separating various types of DNA. It is a reagent in analytical chemistry, where it is used to identify ions by the color and morphology of the precipitate. When enriched in radioisotopes, such as ¹³⁷CsCl or ¹³¹CsCl, caesium chloride is used in nuclear medicine applications such as treatment of cancer and diagnosis of myocardial infarction. Another form of cancer treatment was studied using conventional non-radioactive CsCl. Whereas conventional caesium chloride has a rather low toxicity to humans and animals, the radioactive form easily contaminates the environment due to the high solubility of CsCl in water. Spread of ¹³⁷CsCl powder from a 93-gram container in 1987 in Goiânia, Brazil, resulted in one of the worst-ever radiation spill accidents killing four, including one child, and directly affecting 249 people.

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.^[8] 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 Fm3m).^[5] The rocksalt structure is also observed at ambient conditions in nanometer-thin CsCl films grown on mica, LiF, KBr and NaCl substrates.^[9]

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.^[10] The crystals are highly hygroscopic and deliquescent. Caesium chloride crystals gradually disintegrate at ambient conditions.^[11] Caesium chloride does not form hydrates.^[12]

In contrast to sodium chloride and potassium chloride, caesium chloride readily dissolves in concentrated hydrochloric acid.^[14][15] Caesium chloride has also a relatively high solubility in formic acid (1077 g/L at 18 °C) and hydrazine; medium solubility in methanol (31.7 g/L at 25 °C) and low solubility in ethanol (7.6 g/L at 25 °C),^{[12][15][16]: 97} 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),^[15] ethylacetate and other complex ethers, butanone, acetophenone, pyridine and chlorobenzene.^[17]

Despite its wide band gap of about 8.35 eV at 80 K,^[2] 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.^[18]

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:^[21]

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,^[22] 2CsCl·CuCl₂, CsCl·2CuCl and CsCl·LiCl,^[23] and with interhalogen compounds:^[24]