Titanium(III) chloride

  (Redirected from Titanium trichloride)

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.

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
Names
Other names
titanium trichloride
titanous chloride
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.028.845 Edit this at Wikidata
EC Number
  • 231-728-9
RTECS number
  • XR1924000
UNII
  • InChI=1S/3ClH.Ti/h3*1H;/q;;;+3/p-3 checkY
    Key: YONPGGFAJWQGJC-UHFFFAOYSA-K checkY
  • InChI=1/3ClH.Ti/h3*1H;/q;;;+3/p-3
    Key: YONPGGFAJWQGJC-DFZHHIFOAS
  • Cl[Ti](Cl)Cl
Properties
TiCl3
Molar mass 154.225 g/mol
Appearance red-violet crystals
hygroscopic
Density 2.64 g/cm3
Melting point 425 °C (797 °F; 698 K) (decomposes)
Boiling point 960 °C (1,760 °F; 1,230 K)
very soluble
Solubility soluble in acetone, acetonitrile, certain amines;
insoluble in ether and hydrocarbons
+1110.0×10−6 cm3/mol
1.4856
Hazards
Main hazards Corrosive
Safety data sheet External MSDS
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 otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is checkY☒N ?)
Infobox references

Structure and bondingEdit

In TiCl3, each titanium atom has one d electron, rendering its derivatives paramagnetic, that is, the substance is attracted into a magnetic field. 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 (α), gamma (γ), and delta (δ). 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 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, the trihalides of the heavier metals hafnium and zirconium engage in metal-metal bonding. 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 reactivityEdit

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] This light-blue complex TiCl3(THF)3 forms when TiCl3 is treated with tetrahydrofuran (THF).[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 blue-colored complex Ti(acac)3.[6]

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:[7]

2 TiCl3 → TiCl2 + TiCl4

The ternary halides, such as A3TiCl6, have structures that depend on the cation (A+) added.[8] Caesium chloride treated with titanium(II) chloride and hexachlorobenzene produces crystalline CsTi2Cl7. In these structures Ti3+ exhibits octahedral coordination geometry.[9]

ApplicationsEdit

TiCl3 is the main Ziegler–Natta catalyst, responsible for most industrial production of polyethylene. The catalytic activities depend strongly on the polymorph of the TiCl3 (α vs. β vs. γ vs. δ) and the method of preparation.[10]

Laboratory useEdit

TiCl3 is also a specialized 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.[11] Titanium trichloride can reduce nitrate to ammonium ion thereby allowing for the sequential analysis of nitrate and ammonia.[12] Slow deterioration occurs in air-exposed titanium trichloride, often resulting in erratic results, such as in reductive coupling reactions.[13]

SafetyEdit

TiCl3 and most of its complexes are typically handled under air-free conditions to prevent reactions with oxygen and moisture. Samples of TiCl3 can be relatively air stable or pyrophoric.[14][15]

ReferencesEdit

  1. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  2. ^ Sherfey, J. M. (2007). "Titanium(III) Chloride and Titanium(III) Bromide". Inorganic Syntheses. Inorganic Syntheses. 6. pp. 57–61. doi:10.1002/9780470132371.ch17. ISBN 9780470132371.
  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 (3): 755–757. doi:10.1021/om060486d.
  4. ^ Handlovic, M.; Miklos, D.; Zikmund, M. (1981). "The structure of trichlorotris(tetrahydrofuran)titanium(III)". Acta Crystallographica B. 37 (4): 811–814. doi:10.1107/S056774088100438X.
  5. ^ Manzer, L. E. (1982). "Tetrahydrofuran Complexes of Selected Early Transition Metals". Inorganic Syntheses. 21: 137. doi:10.1002/9780470132524.ch31.
  6. ^ Arslan, Evrim; Lalancette, Roger A.; Bernal, Ivan (2017). "An Historic and Scientific Study of the Poperties of Metal(III) Tris-acetylacetonates". Structural Chemistry. 28: 201–212. doi:10.1007/s11224-016-0864-0. S2CID 99668641.
  7. ^ Holleman, A. F.; Wiberg, E. (2001). Inorganic Chemistry. San Diego, CA: Academic Press. ISBN 0-12-352651-5.[page needed]
  8. ^ Hinz, D.; Gloger, T.; 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.
  9. ^ Jongen, L.; Meyer, G. (2004). "Caesium heptaiododititanate(III), CsTi2I7". Zeitschrift für Anorganische und Allgemeine Chemie. 630 (2): 211–212. doi:10.1002/zaac.200300315.
  10. ^ Whiteley, Kenneth S.; Heggs, T. Geoffrey; Koch; Mawer, Ralph L.; Immel, Wolfgang (2005). "Polyolefins". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a21_487.
  11. ^ Gundersen, Lise-Lotte; Rise, Frode; Undheim, Kjell; Méndez Andino, José. "Titanium(III) Chloride". Encyclopedia of Reagents for Organic Synthesis. doi:10.1002/047084289X.rt120.pub2.
  12. ^ Rich, D. W.; Grigg, B.; Snyder, G. H. (2006). "Determining Ammonium & Nitrate ions using a Gas Sensing Ammonia Electrode". Soil and Crop Science Society of Florida. 65.
  13. ^ Fleming, Michael P.; McMurry, John E. (1981). "Reductive Coupling of Carbonyls to Alkenes: Adamantylideneadamantane". Organic Syntheses. 60: 113. doi:10.15227/orgsyn.060.0113.
  14. ^ Ingraham, T. R.; Downes, K. W.; Marier, P. (1957). "The Production of Titanium Trichloride by Arc-Induced Hydrogen Reduction of Titanium Tetrachloride". Canadian Journal of Chemistry. 35 (8): 850–872. doi:10.1139/v57-118. ISSN 0008-4042.
  15. ^ Pohanish, Richard P.; Greene, Stanley A. (2009). Wiley Guide to Chemical Incompatibilities (3rd ed.). John Wiley & Sons. p. 1010. ISBN 9780470523308.