Interhalogen

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An interhalogen compound is a molecule whose atom contains two or more different halogen atoms (fluorine, chlorine, bromine, iodine, or astatine).

Most interhalogen compounds known are binary (composed of only two distinct elements). Their formulae are generally XYn, where n = 1, 3, 5 or 7, and X is the less electronegative of the two halogens. They are all prone to hydrolysis, and ionise to give rise to polyhalogen ions.

No interhalogen compounds containing three or more different halogens are definitely known,[1] although a few books claim that IFCl
2
and IF
2
Cl
have been obtained,[2][3][4][5] and theoretical studies seem to indicate that some compounds in the series BrClF
n
are barely stable.[6]

Types of interhalogens[edit]

Diatomic interhalogens[edit]

The interhalogens of form XY have physical properties intermediate between those of the two parent halogens. The covalent bond between the two atoms has some ionic character, the less electronegative element, X, being oxidised and having a partial positive charge. All combinations of F, Cl, Br and I which have the above-mentioned general formula are known, but not all are stable, and some combinations of At with other halogens are not even known. Those which are known are highly unstable.

  • Bromine monofluoride (BrF) has not been obtained as a pure compound — it dissociates into the trifluoride and free bromine. It is created according to the following equation:
Br2(l) + F2(g) → 2BrF(g)
  • Iodine monochloride (ICl) exists as red transparent crystals which melt at 27.2 °C to form a choking brownish liquid (similar in appearance and weight to bromine). It reacts with HCl to form the strong acid HICl2. The crystal structure of iodine monochloride consists of puckered zig-zag chains, with strong interactions between the chains.
  • Astatine monochloride (AtCl) is made either by the direct combination of gas-phase astatine with chlorine or by the sequential addition of astatine and dichromate ion to an acidic chloride solution.
  • Iodine monobromide (IBr) is made by the direct combination of the elements to form a dark red crystalline solid. It melts at 42 °C and boils at 116 °C to form a partially dissociated vapour.
  • Astatine monobromide (AtBr) is made by the direct combination of astatine with either bromine vapour or an aqueous solution of iodine monobromide.
  • Astatine monoiodide (AtI) is made by direct combination of astatine and iodine. It is the heaviest confirmed interhalogen compound.

Tetratomic interhalogens[edit]

  • Chlorine trifluoride (ClF3) is a colourless gas which condenses to a green liquid, and freezes to a white solid. It is made by reacting chlorine with an excess of fluorine at 250 °C in a nickel tube. It reacts more violently than fluorine, often explosively. The molecule is planar and T-shaped. It is used in the manufacture of uranium hexafluoride.
  • Bromine trifluoride (BrF3) is a yellow-green liquid which conducts electricity — it ionises to form [BrF2]+ + [BrF4]>. It reacts with many metals and metal oxides to form similar ionised entities; with some others it forms the metal fluoride plus free bromine and oxygen. It is used in organic chemistry as a fluorinating agent. It has the same molecular shape as chlorine trifluoride.
  • Iodine trifluoride (IF3) is a yellow solid which decomposes above −28 °C. It can be synthesised from the elements, but care must be taken to avoid the formation of IF5. F2 attacks I2 to yield IF3 at −45 °C in CCl3F. Alternatively, at low temperatures, the fluorination reaction I2 + 3XeF2 → 2IF3 + 3Xe can be used. Not much is known about iodine trifluoride as it is so unstable.
  • Iodine trichloride (ICl3) forms lemon yellow crystals which can be melted under pressure to a brown liquid. It can be made from the elements at low temperature, or from iodine pentoxide and hydrogen chloride. It reacts with many metal chlorides to form tetrachloroiodides, and hydrolyses in water. The molecule is a planar dimer (ICl3)2, with each iodine atom surrounded by four chlorine atoms.

Hexatomic interhalogens[edit]

  • Chlorine pentafluoride (ClF5) is a colourless gas, made by reacting chlorine trifluoride with fluorine at high temperatures and high pressures. It reacts violently with water and most metals and nonmetals.
  • Bromine pentafluoride (BrF5) is a colourless fuming liquid, made by reacting bromine trifluoride with fluorine at 200 °C. It is physically stable, but reacts violently with water and most metals and nonmetals.
  • Iodine pentafluoride (IF5) is a colourless liquid, made by reacting iodine pentoxide with fluorine, or iodine with silver(II) fluoride. It is highly reactive, even slowly with glass. It reacts with elements, oxides and carbon halides. The molecule has the form of a tetragonal pyramid.
  • Iodine pentabromide (IBr5), if it exists (there is some dispute on this point),[7] is a dark reddish-brown liquid or brown-yellow to colorless crystalline solid, made by reacting iodine with bromine at 60 °C. In its liquid state it resembles bromine in most properties; in any state, it is very toxic. It is unstable upon heating above the boiling point of bromine, yielding bromine vapor and iodine monobromide.[8][9][10]

Octatomic interhalogens[edit]

  • Iodine heptafluoride (IF7) is a colourless gas and a strong fluorinating agent. It is made by reacting iodine pentafluoride with fluorine gas. The molecule is a pentagonal bipyramid. This compound is the only known interhalogen compound where the larger atom is carrying seven of the smaller atoms.
  • All attempts to synthesize bromine heptafluoride have met with failure; instead, bromine pentafluoride and fluorine gas are produced.

Summary of known interhalogens[edit]

F Cl Br I At
F
F2
Cl
ClF, ClF3, ClF5
Cl2
Br
BrF, BrF3, BrF5
BrCl
Br2
I
IF, IF3, IF5, IF7
ICl, (ICl3)2
IBr, IBr5(?)
I2
At
AtCl
AtBr
AtI
At2(?)

Properties[edit]

Typically, interhalogen bonds are more reactive than diatomic halogen bonds because interhalogen bonds are weaker than diatomic halogen bonds except for F2. If interhalogens are exposed to water, they will convert to halide and oxyhalide ions. With BrF5, this reaction can be explosive. However, if interhalogens are exposed to silicon dioxide, or metal oxides, then silicon or metal respectively will bond with one of the types of halogen, leaving free diatomic halogens and diatomic oxygen. Most interhalogens are halogen fluorides, and all but three (IBr, AtBr, and AtI) of the remainder are halogen chlorides. Chlorine and bromine can each bond to five fluorine atoms, and iodine can bond to seven. AX and AX3 interhalogens can form between two halogens whose electronegativities are relatively close to one another. When interhalogens are exposed to metals, they react to form metal halides of the constituent halogens. The oxidization power of an interhalogen increases with the number of halogens attached to the central atom of the interhalogen, as well as with the decreasing size of the central atom of the compound. Interhalogens containing fluorine are more likely to be volatile than interhalogens containing heavier halogens.[1]

Interhalogens with one or three halogens bonded to a central atom are formed by two elements whose electronegativities are not far apart. Interhalogens with five or seven halogens bonded to a central atom are formed by two elements whose sizes are very different. The number of smaller halogens that can bond to a large central halogen is guided by the ratio of the atomic radius of the larger halogen over the atomic radius of the smaller halogen. A number of interhalogens, such as IF7, react with all metals except for those in the platinum group. IF7, unlike interhalogens in the XY5 series, does not react with the fluorides of the alkali metals.[1]

ClF3 is the most reactive of the XY3 interhalogens. ICl3 is the least reactive. BrF3 has the highest thermal stability of the interhalogens with four atoms. ICl3 has the lowest. Chlorine trifluoride has a boiling point of −12° Celsius. Bromine trifluoride has a boiling point of 127° Celsius and is a liquid at room temperature. Iodine trichloride melts at 101° Celsius.[1]

Most interhalogens are covalent gases. However, some interhalogens, especially those containing bromine, are liquids, and most iodine-containing interhalogens are solids. Most of the interhalogens composed of lighter halogens are fairly colorless, but the interhalogens containing heavier halogens are deeper in color due to their higher molecular weight. In this respect, the interhalogens are similar to the halogens. The greater the difference between the electronegativities of the two halogens in an interhalogen, the higher the boiling point of the interhalogen. All interhalogens are diamagnetic. The bond length of interhalogens in the XY series increases with the size of the constituent halogens. For instance, ClF has a bond length of 1.628 angstroms, and IBr has a bond length of 2.47 angstroms.[1]

Production[edit]

It is possible to produce larger interhalogens, such as ClF3, by exposing smaller interhalogens, such as ClF, to pure diatomic halogens, such as F2. This method of production is especially useful for generating halogen fluorides. However, at temperatures of 250 to 300° Celsius, this type of production method can also convert larger interhalogens into smaller ones. It is also possible to produce interhalogens by combining two pure halogens at various conditions. This method can generate any interhalogen save for IF7.[1]

Smaller interhalogens, such as ClF, can be formed by direct reaction with pure halogens. For instance, F2 will react with Cl2 at 250° Celsius to form two molecules of ClF. Br2 will react with diatomic fluorine in the same way, but at 60 °C. I2 reacts with diatomic fluorine at only 35° Celsius. ClF and BrF can both be produced by the reaction of a larger interhalogen, such as ClF3 or BrF3 and a diatomic molecule of the element lower in the periodic table. Among the hexatomic interhalogens, IF5 has a higher boiling point (97° Celsius) than BrF5 (40.5° Celsius), although both compounds are liquids at room temperature. The interhalogen IF7 can be formed by reacting palladium iodide with fluorine.[1]

Applications[edit]

Some interhalogens, such as BrF3, IF5, and ICl, are good halogenating agents. However, BrF5 is too reactive to generate fluorine. Beyond that, iodine monochloride has several applications, including helping to measure the saturation of fats and oils, and as a catalyst for some reactions. A number of interhalogens, including IF7, are used to form polyhalides.[1]

See also[edit]

References[edit]

  1. ^ a b c d e f g h P.B. Saxena (2007), Chemistry Of Interhalogen Compounds, retrieved February 27, 2013 
  2. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 824. ISBN 0080379419. 
  3. ^ Robert A. Meyers, editor (2001), "Encyclopedia of Physical Science and Technology: Inorganic Chemistry", third edition. Academic Press. ISBN 978-0-12-227410-7 Quote: "A few ternary compounds, such as IFCl
    2
    and IF
    2
    Cl
    , are also known.
    " (no source given)
  4. ^ C. Parameshwara Murthy (2008), "University Chemistry", volume 1, 675 pages. New Age International. ISBN 8122407420. Quote: "The only two interhalogen compounds are IFCl
    2
    and IF
    2
    Cl
    " (no source given)
  5. ^ Balaram Sahoo, Nimai Charan Nayak, Asutosh Samantaray, Prafulla Kumar Pujapanda (2012), "Inorganic Chemistry". PHI Learning Pvt. Ltd. ISBN 8120343085. Quote: "Only a few ternary interhalogen compounds such as IFCl
    2
    and IF
    2
    Cl
    have been preprared.
    " (no source given)
  6. ^ Igor S. Ignatyev and Henry F. Schaefer III (1999), "Bromine Halides:  The Neutral Molecules BrClF
    n
    (n = 1-5) and Their Anions  Structures, Energetics, and Electron Affinities". Journal of the American Chemical Society, volume 121, issue 29, pages 6904–6910. doi:10.1021/ja990144h Conclusion: maybe there are some barely stable compounds.
  7. ^ Sharpe (1956). Supplement to Mellor's comprehensive treatise on inorganic and theoretical chemistry, Supplement II, Part 1, (F, Cl, Br, I, At). p. 742. 
  8. ^ Hare, Hobart Amory; Caspari, Charles; Rusby, Henry Hurd; Geisler, Joseph Frank; Kremers, Edward; Base, Daniel (1909). The national standard dispensatory (2nd ed.). Lea & Febiger. pp. 858–859. 
  9. ^ Manahan, Stanley E (2003). Toxicological Chemistry and Biochemistry (3rd ed.). CRC Press. p. 241. ISBN 9781566706186. 
  10. ^ Grushko, Ya. M (1992). Handbook of Dangerous Properties of Inorganic And Organic Substances in Industrial Wastes. CRC Press. p. 54. ISBN 9780849393006. 

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