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Chromium(III) chloride (also called chromic chloride) is a violet coloured solid with the formula CrCl3.
Although it is ionic, the solid state structure is kinetically inert so that anhydrous CrCl3 is surprisingly reluctant to dissolve in water. However, in the presence of a trace of a reducing agent capable of reducing Cr3+ to Cr2+, the CrCl3 dissolves rapidly to form soluble complexes containing hydrated Cr3+ ions. The common commercial form of the hydrate is the dark green complex shown in the picture, [CrCl2(H2O)4]Cl.2H2O, but two other forms are known, viz., pale green [CrCl(H2O)5]Cl2.H2O and violet [Cr(H2O)6]Cl3.
This inertness means that CrCl3 is generally sluggish to react without the presence of a reducing agent. When it does react it undergoes ligand substitution reactions to form other complexes of chromium(III). It reacts as a Lewis acid, forming stable chloro complexes such as [CrCl6]3-.
Chromium(III) chloride is a Lewis acid, classified as "hard" according to the Hard-Soft Acid-Base theory. However it is also a chloro complex which is quite inert to substitution, so in fact it is ordinarily quite unreactive. 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.
The most common form of CrCl3 sold commercially is a dark green hexahydrate with the structure [CrCl2(H2O)4]Cl.2H2O, and like the anhydrous form it is also very inert towards substitution.
If substitution reactions are performed in the presence of a trace of Cr2+, then CrCl3 can undergo substitution with ligands such as water (giving violet [Cr(H2O)6]3+) or pyridine:
CrCl3 + 3 C5H5N → [CrCl3(C5H5N)3]
Such complexes are usually octahedral.
With molten alkali metal chlorides such as potassium chloride, CrCl3 gives octahedral complexes of the type K3CrCl6, as well as K3Cr2Cl9 which is also octahedral but where the two chromiums are linked via three chloride bridges.
It may also be prepared from the hexahydrate, by heating with thionyl chloride which reacts with the water of hydration.
A significant use of CrCl3 in organic synthesis is for the in situ preparation of chromium(II) chloride, a popular reagent for (A) reduction of alkyl halides and for (B) the synthesis of (E)-alkenyl halides. The reaction is usually performed using two moles of CrCl3 per mole of lithium aluminium hydride, although if aqueous acidic conditions are appropriate zinc and hydrochloric acid may be sufficient.
Chromium(III) chloride has also been used as a Lewis acid in organic reactions, for example to catalyse the nitroso Diels-Alder reaction.8 it has a tae
Although trivalent chromium is far less poisonous than hexavalent, chromium salts are generally considered highly toxic. Avoid ingestion and inhalation of dust. Wear gloves and goggles.
|This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Chromium(III)_chloride". A list of authors is available in Wikipedia.|