Oxohalide

In chemistry, molecular oxohalides (oxyhalides) are a group of chemical compounds in which both oxygen and halogen atoms are attached to another chemical element A in a single molecule. They have the general formula AO_mX_n, where X is a halogen. Known oxohalides have fluorine (F), chlorine (Cl), bromine (Br), and/or iodine (I) in their molecules. The element A may be a main group element, a transition element, a rare earth element or an actinide. The term oxohalide, or oxyhalide, may also refer to minerals and other crystalline substances with the same overall chemical formula, but having an ionic structure.

Oxohalides can be seen as compounds intermediate between oxides and halides. There are three general methods of synthesis:^[1]

• Partial oxidation of a halide:

  2 PCl₃ + O₂ → 2 POCl₃

  • In this example, the oxidation state increases by two and the electrical charge is unchanged.
• Partial halogenation of an oxide:

  2 V₂O₅ + 6 Cl₂ + 3 C → 4 VOCl₃ + 3 CO₂

• Oxide replacement:

  CrO2−4 + 2 Cl⁻ + 4 H⁺ → CrO₂Cl₂ + 4 H₂O

In addition, various oxohalides can be made by halogen exchange reactions and this reaction can also lead to the formation of mixed oxohalides such as POFCl₂ and CrO₂FCl.

In relation to the oxide or halide, for a given oxidation state of an element A, if two halogen atoms replace one oxygen atom, or vice versa, the overall charge on the molecule is unchanged and the coordination number of the central atom decreases by one. For example, both phosphorus oxychloride (POCl₃) and phosphorus pentachloride, (PCl₅) are neutral covalent compounds of phosphorus in the +5 oxidation state. If an oxygen atom is simply replaced by a halogen atom the charge increases by +1, but the coordination number is unchanged. This is illustrated by the reaction of a mixture of a chromate or dichromate salt and potassium chloride with concentrated sulfuric acid.

Cr₂O2−7 + 4 Cl⁻ + 6 H⁺ → 2 CrO₂Cl₂ + 3 H₂O

The chromyl chloride produced has no electrical charge and is a volatile covalent molecule that can be distilled out of the reaction mixture.^[2]

Oxohalides of elements in high oxidation states are strong oxidizing agents, with oxidizing power similar to the corresponding oxide or halide. Most oxohalides are easily hydrolyzed. For example, chromyl chloride is hydrolyzed to chromate in the reverse of the synthetic reaction, above. The driving force for this reaction is the formation of A-O bonds which are stronger than A-Cl bonds. This gives a favourable enthalpy contribution to the Gibbs free energy change for the reaction^[3]

Many oxohalides can act as Lewis acids. This is particularly so with oxohalides of coordination number 3 or 4 which, in accepting one or more electron pairs from a Lewis base, become 5- or 6-coordinate. Oxohalide anions such as [VOCl₄]²⁻ can be seen as acid-base complexes of the oxohalide (VOCl₂) with more halide ions acting as Lewis bases. Another example is VOCl₂ which forms the trigonal bipyramidal complex VOCl₂(N(CH₃)₃)₂ with the base trimethylamine.^[4]

The vibrational spectra of many oxohalides have been assigned in detail. They give useful information on relative bond strengths. For example, in CrO₂F₂, the Cr–O stretching vibrations are at 1006 cm⁻¹ and 1016 cm⁻¹ and the Cr–F stretching vibrations are at 727 cm⁻¹ and 789 cm⁻¹. The difference is much too large to be due to the different masses of O and F atoms. Rather, it shows that the Cr–O bond is much stronger than the Cr–F bond. M–O bonds are generally considered to be double bonds and this is backed up by measurements of M–O bond lengths. It implies that the elements A and O are chemically bound together by a σ bond and a π bond.^[5]

Oxohalides of elements in high oxidation states are intensely coloured owing to ligand to metal charge transfer (LMCT) transitions.^[6]

Carbon forms oxohalides COX₂, X = F, Br, and the very toxic phosgene (X = Cl), which is produced industrially by a carbon-catalyzed reaction of carbon monoxide with chlorine. It is a useful reagent in organic chemistry for the formation of carbonyl compounds.^[7] For example:

COCl₂ + 2 ROH → CO(OR)₂ + 2 HCl

Silicon tetrafluoride reacts with water to yield poorly-characterized oxyfluoride polymers, but slow and careful reaction at -196 °C yields the oxyfluoride hexafluorodisiloxane as well.^[8]

Nitrogen forms two series of oxohalides with nitrogen in oxidation states 3, NOX, X = F, Cl, Br and 5, NO₂X, X = F, Cl. They are made by halogenation of nitrogen oxides. Note that NO₂F is isoelectronic with the nitrate ion, NO−3. Only oxohalides of phosphorus(V) are known.^[9]

Sulfur forms oxohalides^[10] in oxidation state +4, such as thionyl chloride, SOCl₂ and oxidation state +6, such as sulfuryl fluoride (SO₂F₂), sulfuryl chloride (SO₂Cl₂), and thionyl tetrafluoride (SOF₄). All are easily hydrolyzed. Indeed, thionyl chloride can be used as a dehydration agent as the water molecules are converted into gaseous products, leaving behind the anhydrous solid chloride.^[11]

MgCl₂·6H₂O + 6 SOCl₂ → MgCl₂ + 6 SO₂ + 12 HCl

Selenium and tellurium form similar compounds and also the oxo-bridged species F₅AOAF₅ (A = S, Se, Te). They are non-linear with the A-O-A angle of 142.5, 142.4 and 145.5° for S, Se and Te, respectively.^[12] The tellurium anion F₅TeO⁻, known as teflate, is a large and rather stable anion, useful for forming stable salts with large cations.^[11]

The halogens form various oxofluorides with formulae XO₂F (chloryl fluoride), XO₃F (perchloryl fluoride) and XOF₃ with X = Cl, Br and I. IO₂F₃ and IOF₅ are also known.^[13]

Xenon forms xenon oxytetrafluoride (XeOF₄), xenon dioxydifluoride (XeO₂F₂) and xenon oxydifluoride (XeOF₂).

A selection of known oxohalides of transition metals is shown below, and more detailed lists are available in the literature.^[15] X indicates various halides, most often F and Cl.

High oxidation states of the metal are dictated by the fact that oxygen is a strong oxidizing agent, as is fluorine. Bromine and iodine are relatively weak oxidizing agents, so fewer oxobromides and oxoiodides are known. Structures for compounds with d⁰ configuration are predicted by VSEPR theory. Thus, CrO₂Cl₂ is tetrahedral, OsO₃F₂ is trigonal bipyramidal, XeOF₄ is square pyramidal and OsOF₅ is octahedral.^[18] The d¹ complex ReOCl₄ is square pyramidal.

The compounds [Ta₂OX₁₀]²⁻ and [M₂OCl₁₀]⁴⁻ (M = W, Ru, Os) have two MX₅ groups joined by a bridging oxygen atom.^[19] Each metal has an octahedral environment. The unusual linear M−O−M structure can be rationalized in terms of molecular orbital theory, indicating the presence of d_π — p_π bonding between the metal and oxygen atoms.^[20] Oxygen bridges are present in more complex configurations like M(cp)₂(OTeF₅)₂ (M = Ti, Zr, Hf, Mo or W; cp = cyclopentadienyl, η⁵-C₅H₅)^[21] or [AgOTeF₅-(C₆H₅CH₃)₂]₂.^[17]

In the actinide series, uranyl compounds such as uranyl chloride (UO₂Cl₂) and [UO₂Cl₄]²⁻ are well known and contain the linear UO₂ moiety. Similar species exist for neptunium and plutonium. The species uranyl fluoride is a complicating contaminant in samples uranium hexafluoride.

Bismuth oxochloride (BiOCl, bismoclite) is a rare example of a mineral oxohalide. The crystal structure has a tetragonal symmetry and can be thought of as consisting of layers of Cl⁻, Bi³⁺ and O²⁻ ions, in the order Cl-Bi-O-Bi-Cl-Cl-Bi-O-Bi-Cl. This layered, graphite-like structure results in a relatively low hardness of bismoclite (Mohs 2–2.5) and most other oxohalide minerals.^[22] Those other minerals include terlinguaite Hg₂OCl, formed by the weathering of mercury-containing minerals.^[23] Mendipite, Pb₃O₂Cl₂, formed from an original deposit of lead sulfide in a number of stages is another example of a secondary oxohalide mineral.

The elements iron, antimony, bismuth and lanthanum form oxochlorides of general formula MOCl. MOBr and MOI are also known for Sb and Bi. Many of their crystal structures have been determined.^[24]

• Transition metal oxo complex

Wikimedia Commons has media related to Oxohalides.

• Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
• Housecroft, C. E. and Sharpe, A. G. Inorganic Chemistry, 2nd ed., Pearson Prentice-Hall 2005. ISBN 0-582-31080-6
• Shrivr, D. F. and Atkins, P. W. Inorganic Chemistry, 3rd edn. Oxford University Press, 1999. ISBN 0-19-850330-X
• Wells, A. F. (1962). Structural Inorganic Chemistry (3rd ed.). Oxford: Clarendon Press. pp. 384–392. ISBN 0-19-855125-8..