Chromium(III) chloride (also called chromic chloride) is an inorganic chemical compound with the chemical formula CrCl₃. This crystalline salt forms several hydrates with the formula CrCl₃·nH₂O, among which are hydrates where n can be 5 (chromium(III) chloride pentahydrate CrCl₃·5H₂O) or 6 (chromium(III) chloride hexahydrate CrCl₃·6H₂O). The anhydrous compound with the formula CrCl₃ are violet crystals, while the most common form of the chromium(III) chloride are the dark green crystals of hexahydrate, CrCl₃·6H₂O. Chromium chlorides find use as catalysts and as precursors to dyes for wool.

Anhydrous chromium(III) chloride adopts the YCl₃ structure, with Cr³⁺ occupying one third of the octahedral interstices in alternating layers of a pseudo-cubic close packed lattice of Cl⁻ ions. The absence of cations in alternate layers leads to weak bonding between adjacent layers. For this reason, crystals of CrCl₃ cleave easily along the planes between layers, which results in the flaky (micaceous) appearance of samples of chromium(III) chloride. The anhydrous CrCl₃ is exfoliable down to the monolayer limit. If pressurized to 9.9 GPa it goes under a phase transition.

The hydrated chromium(III) chlorides display the somewhat unusual property of existing in a number of distinct chemical forms (isomers), which differ in terms of the number of chloride anions that are coordinated to Cr(III) and the water of crystallization. The different forms exist both as solids and in aqueous solutions. Several members are known of the series of [CrCl_3−q(H₂O)_n]^q+. The common hexahydrate can be more precisely described as [CrCl₂(H₂O)₄]Cl·2H₂O. It consists of the cation trans-[CrCl₂(H₂O)₄]⁺ and additional molecules of water and a chloride anion in the lattice. Two other hydrates are known, pale green [CrCl(H₂O)₅]Cl₂·H₂O and violet [Cr(H₂O)₆]Cl₃. Similar hydration isomerism is seen with other chromium(III) compounds.^{[citation needed]}

Anhydrous chromium(III) chloride may be prepared by chlorination of chromium metal directly, or indirectly by carbothermic chlorination of chromium(III) oxide at 650–800 °C

The hydrated chlorides are prepared by treatment of chromate with hydrochloric acid and aqueous methanol.

Slow reaction rates are common with chromium(III) complexes. The low reactivity of the d³ Cr³⁺ ion can be explained using crystal field theory. One way of opening CrCl₃ up to substitution in solution is to reduce even a trace amount to CrCl₂, for example using zinc in hydrochloric acid. This chromium(II) compound undergoes substitution easily, and it can exchange electrons with CrCl₃ via a chloride bridge, allowing all of the CrCl₃ to react quickly. With the presence of some chromium(II), solid CrCl₃ dissolves rapidly in water. Similarly, ligand substitution reactions of solutions of [CrCl₂(H₂O)₄]⁺ are accelerated by chromium(II) catalysts.

With molten alkali metal chlorides such as potassium chloride, CrCl₃ gives salts of the type M₃[CrCl₆] and K₃[Cr₂Cl₉], which is also octahedral but where the two chromiums are linked via three chloride bridges.

The hexahydrate can also be dehydrated with thionyl chloride: