Titanium(III) chloride

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Titanium(III) chloride
Beta-TiCl3-chain-from-xtal-3D-balls.png
Beta-TiCl3-chains-packing-from-xtal-3D-balls-B.png
β-TiCl3 viewed along the chains
TiCl3.jpg
TiCl3 solution
Identifiers
CAS number 7705-07-9 YesY
PubChem 62646
ChemSpider 56398 YesY
EC number 231-728-9
RTECS number XR1924000
Jmol-3D images Image 1
Properties
Molecular formula TiCl3
Molar mass 154.225 g/mol
Appearance red-violet crystals
hygroscopic
Density 2.64 g/cm3
Melting point 425 °C (decomposes)
Boiling point 960 °C
Solubility in water very soluble
Solubility soluble in acetone, acetonitrile, certain amines;
insoluble in ether and hydrocarbons
Refractive index (nD) 1.4856
Hazards
MSDS External MSDS[dead link]
EU Index Not listed
Main hazards Corrosive
Related compounds
Other anions Titanium(III) fluoride
Titanium(III) bromide
Titanium(III) iodide
Other cations Scandium(III) chloride
Chromium(III) chloride
Vanadium(III) chloride
Related compounds Titanium(IV) chloride
Titanium(II) chloride
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
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Infobox references

Titanium(III) chloride is the inorganic compound with the formula TiCl3. At least four distinct species have this formula; additionally hydrated derivatives are known. TiCl3 is one of the most common halides of titanium and is an important catalyst for the manufacture of polyolefins.

Structure and bonding[edit]

In TiCl3, each Ti atom has one d electron, rendering its derivatives paramagnetic, i.e. the substance is attracted into a magnetic field. The trihalides of hafnium and zirconium: in these heavier metals engage in metal-metal bonding. Solutions of titanium(III) chloride are violet, which arises from excitations of its d-electron. The colour is not very intense since the transition is forbidden by the Laporte selection rule.

Four solid forms or polymorphs of TiCl3 are known. All feature titanium in an octahedral coordination sphere. These forms can be distinguished by crystallography as well as by their magnetic properties, which probes exchange interactions. β-TiCl3 crystallizes as brown needles. Its structure consists of chains of TiCl6 octahedra that share opposite faces such that the closest Ti—Ti contact is 2.91 Å. This short distance indicates strong metal-metal interactions (See Figure in upper right). The three violet "layered" forms, named for their color and their tendency to flake, are called alpha, gamma, and delta. In α-TiCl3, the chloride anions are hexagonal close-packed. In γ-TiCl3, the chlorides anions are cubic close-packed. Finally, disorder in shift successions, causes an intermediate between alpha and gamma structures, called the delta (δ) form. The TiCl6 share edges in each form, with 3.60 Å being the shortest distance between the titanium cations. This large distance between titanium cations precludes direct metal-metal bonding. In contrast, direct Zr-Zr bonding is indicated in zirconium(III) chloride. The difference between the Zr(III) and Ti(III) materials is attributed in part to the relative radii of these metal centers.[1]

Synthesis and reactivity[edit]

TiCl3 is produced usually by reduction of titanium(IV) chloride. Older reduction methods used hydrogen:[2]

2 TiCl4 + H2 → 2 HCl + 2 TiCl3

It is conveniently reduced with aluminium and sold as a mixture with aluminium trichloride, TiCl3·AlCl3. This mixture can be separated to afford TiCl3(THF)3.[3] The complex adopts a meridional structure.[4]

Its hydrate can be synthesised by dissolving titanium in aqueous hydrochloric acid.

2 Ti + 6 HCl + 3 H2O → 2 TiCl3(H2O)3 + 3 H2

TiCl3 forms a variety of coordination complexes, most of which are octahedral. The light-blue crystalline adduct TiCl3(THF)3 forms when TiCl3 is treated with tetrahydrofuran.[5]

TiCl3 + 3 C4H8O → TiCl3(OC4H8)3

An analogous dark green complex arises from complexation with dimethylamine. In a reaction where all ligands are exchanged, TiCl3 is a precursor to the tris acetylacetonate complex.

The more reduced titanium(II) chloride is prepared by the thermal disproportionation of TiCl3 at 500 °C. The reaction is driven by the loss of volatile TiCl4:[6]

2 TiCl3 → TiCl2 + TiCl4

The ternary halides, such as A3TiCl6, have structures that depend on the cation (A+) added.[7] Caesium chloride treated with titanium(II) chloride and hexachlorobenzene produces crystalline CsTi2Cl7. In these structures Ti3+ exhibits octahedral coordination geometry.[8]

Applications[edit]

TiCl3 is the main Ziegler-Natta catalyst, responsible for most industrial production of polypropylene. The catalytic activities depend strongly on the polymorph and the method of preparation.[9]

Laboratory use[edit]

TiCl3 is also a reagent in organic synthesis, useful for reductive coupling reactions, often in the presence of added reducing agents such as zinc. It reduces oximes to imines.[10] Titanium trichloride can reduce nitrate to ammonium ion thereby allowing for the sequential analysis of nitrate and ammonia.[11] Slow deterioration occurs in air-exposed titanium trichloride, often resulting in erratic results, e.g. in reductive coupling reactions.[12]

Safety[edit]

TiCl3 and most of its complexes are typically handled under an air-free conditions to prevent reactions with oxygen and moisture. Depending on the method for its preparation, samples of TiCl3 can be relatively air stable or pyrophoric.[2][13]

References[edit]

  1. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0080379419. 
  2. ^ a b T. R. Ingraham, K. W. Downes, P. Marier, "Titanium(III) Chloride" Inorganic Syntheses, 1960, vol 6, pp. 52–56. doi: 10.1002/9780470132371.ch16
  3. ^ Jones, N. A.; Liddle, S. T.; Wilson, C.; Arnold, P. L. (2007). "Titanium(III) Alkoxy-N-heterocyclic Carbenes and a Safe, Low-Cost Route to TiCl3(THF)3". Organometallics 26: 755–757. doi:10.1021/om060486d. 
  4. ^ Handlovic, M.; Miklos, D.; Zikmund, M. "The structure of trichlorotris(tetrahydrofuran)titanium(III)" Acta Crystallographica 1981, volume B37(4), 811-14.doi:10.1107/S056774088100438X
  5. ^ Manzer, L. E.; Deaton, Joe; Sharp, Paul; Schrock, R. R. (1982). "Tetrahydrofuran Complexes of Selected Early Transition Metals". Inorg. Synth. 21: 137. doi:10.1002/9780470132524.ch31. 
  6. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  7. ^ Hinz, D.; Gloger, T. and Meyer, G. (2000). "Ternary halides of the type A3MX6. Part 9. Crystal structures of Na3TiCl6 and K3TiCl6". Zeitschrift für Anorganische und Allgemeine Chemie 626 (4): 822–824. doi:10.1002/(SICI)1521-3749(200004)626:4<822::AID-ZAAC822>3.0.CO;2-6. 
  8. ^ Jongen, L. and Meyer, G. (2004). "Caesium heptaiododititanate(III), CsTi2I7". Zeitschrift für Anorganische und Allgemeine Chemie 630 (2): 211–212. doi:10.1002/zaac.200300315. 
  9. ^ Kenneth S. Whiteley,T. Geoffrey Heggs, Hartmut Koch, Ralph L. Mawer, Wolfgang Immel, "Polyolefins" in Ullmann's Encyclopedia of Industrial Chemistry 2005, Wiley-VCH, Weinheim. doi:10.1002/14356007.a21_487
  10. ^ Lise-Lotte Gundersen, Frode Rise, Kjell Undheim, José Méndez-Andino, "Titanium(III) Chloride" in Encyclopedia of Reagents for Organic Synthesis doi:10.1002/047084289X.rt120.pub2
  11. ^ "Determining Ammonium & Nitrate ions using a Gas Sensing Ammonia Electrode". Soil and Crop Science Society of Florida, Vol. 65, 2006, D.W.Rich, B.Grigg, G.H.Snyder
  12. ^ Fleming, M. P; McMurry, J. E., "Reductive Coupling of Carbonyls to Alkenes: Adamantylideneadamantane", Org. Synth. ; Coll. Vol. 7: 1 
  13. ^ Pohanish, Richard P. and Greene, Stanley A. (2009). Wiley Guide to Chemical Incompatibilities (3 ed.). John Wiley & Sons. p. 1010. ISBN 9780470523308.