Chloryl

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Chloryl
Structural formula
Space-filling model of crystal structure
Names
Preferred IUPAC name
Chloryl
Systematic IUPAC name
Dioxo-λ5-chloranylium
Identifiers
3D model (JSmol)
ChemSpider
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Infobox references

In chemistry, chloryl refers to a triatomic cation with chemical formula ClO+
2
. This species has the same general structure as chlorite (ClO
2
) but it is electronically different, with chlorine having a +5 oxidation state (rather than the +3 of chlorite). This makes it a rare example of a positively charged oxychloride. Chloryl compounds, such as FClO
2
and [ClO2][RuF6], are all highly reactive and react violently with water and most organic compounds.[1][2]

Structure[edit]

The ClO+
2
cation is isoelectronic with SO
2
,[3] and has a bent structure with a bond angle close to 120°. The Cl–O bond is of bond order 1.5, with its Lewis structure consisting of a double bond and a dative bond which does not utilize d-orbitals.[4]

The red color of ClO+
2
is caused by electron transitions into an antibonding orbital. The analogous transition in SO
2
is not in the visible spectrum, so SO
2
is colorless. The strength of interaction with the counterion affects the energy of this antibonding orbital; thus, in colorless chloryl compounds, strong interactions with the counterion, corresponding with the higher covalent character of the bonding, shift the transition energy out of the visible spectrum.[3]

Compounds[edit]

Thermal ellipsoid model of the coordination environment of chlorine in chloryl hexafluoroantimonate, [ClO2][SbF6], showing F-Cl interactions.

There are two categories of chloryl compounds. The first category is colorless, and includes chloryl fluoride (FClO
2
). These are moderately reactive. Although named as an ionic "chloryl" compound, chloryl fluoride is more a covalent compound than an ionic compound of fluoride and chloryl cation.

The second category features red-colored compounds that are highly reactive. These include chloryl fluorosulfate, ClO
2
SO
3
F
, and dichloryl trisulfate, (ClO
2
)
2
(S
3
O
10
)
. These chloryl compounds form red solutions in fluorosulfuric acid, and do contain a red-colored ClO+
2
cation which dissociates in solution. In the solid state, the Raman and infrared spectra indicate strong interactions with the counterion.[3][1] Not all chloryl compounds in the solid state are necessarily ionic. The reaction products of FClO
2
with BF
3
and PF
5
are assumed to be molecular adducts rather than true salts.[3][5]

One notable chloryl compound is dichlorine hexoxide, which exists as an ionic compound more accurately described as chloryl perchlorate, [ClO
2
]+
[ClO
4
]
.[6] It is a red fuming liquid under standard conditions.

Chloryl compounds are best prepared by the reaction of FClO
2
with a strong Lewis acid. For example:[5]

FClO
2
+ AsF
5
→ [ClO2][AsF6]

Other synthesis routes are also possible, including:[5]

5 ClO
2
+ 3 AsF
5
→ 2 [ClO2][AsF6] + AsF
3
O
+ 4 Cl
2
Cl
2
O
6
+ 2 SbF
5
→ [ClO2][SbF6] + SbF
3
O
+ FClO
3

Metathesis reactions may be carried out with strong Lewis bases. For example, the reaction of the hexafluoroplatinate salt with nitryl fluoride yields the nitronium salt:[5]

[ClO2][PtF6] + FNO
2
→ [NO2][PtF6] + FClO
2

References[edit]

  1. ^ a b Christe, K. O.; Schack, C. J.; Pilipovich, D.; Sawodny, W. (1969). "Chloryl cation, ClO+
    2
    ". Inorganic Chemistry. 8 (11): 2489–2494. doi:10.1021/ic50081a050.
     
  2. ^ Bougon, R.; Cicha, W. V.; Lance, M.; Meublat, L.; Nierlich, M.; Vigner, J. (1991). "Preparation characterization and crystal structure of chloryl hexafluororuthenate(1-). Crystal structure of [ClF
    2
    ]+
    [RuF
    6
    ]
    ". Inorganic Chemistry. 30 (1): 102–109. doi:10.1021/ic00001a019.
     
  3. ^ a b c d Carter, H. A.; Johnson, W. M.; Aubke, F. (15 December 1969). "Chloryl compounds. Part II. Chloryl hexafluoroarsenate and chloryl fluoride". Canadian Journal of Chemistry. 47 (24): 4619–4625. doi:10.1139/v69-763. 
  4. ^ David L. Cooper (2001). "Spin-coupled description of the chemical bonding to hypercoordinate chlorine". Theoretical Chemistry Accounts. 105 (4-5): 323–327. doi:10.1007/PL00013292. 
  5. ^ a b c d K. O. Christe; C. J. Schack (1976). Harry Julius Emeléus, A. G. Sharpe, ed. Chlorine Oxyfluorides. Advances in inorganic chemistry and radiochemistry, Volume 18. Academic Press. pp. 356–358. ISBN 0-12-023618-4. 
  6. ^ Tobias, K. M.; Jansen, M. (1986). "Crystal Structure of Cl
    2
    O
    6
    ". Angewandte Chemie International Edition in English. 25 (11): 993–994. doi:10.1002/anie.198609931.