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Titanium tetrachloride

Safety data
IUPAC name Titanium tetrachloride
Titanium(IV) chloride
CAS number 7550-45-0
Molecular formula TiCl4
Molar mass 189.71 g/mol
Appearance colourless fuming liquid
Density 1.730 g/ml, liquid
Melting point

-24 °C

Boiling point

136.4 °C

Solubility in water Decomposes
Molecular shape Tetrahedral
Dipole moment zero
Std enthalpy of
-804.16 kJ/mol
Standard molar
221.93 J·K−1·mol−1
EU classification Corrosive
NFPA 704
R-phrases R14, R34
S-phrases (S1)/(S2), S7/S8, S26,
S36/S37/S39, S45
Related Compounds
Other anions Titanium(IV) fluoride
Titanium(IV) bromide
Titanium(IV) iodide
Other cations Zirconium(IV) chloride
Hafnium(IV) chloride
Related compounds Titanium(II) chloride
Titanium(III) chloride
Except where noted otherwise, data are given for
materials in their standard state
(at 25 °C, 100 kPa)

Infobox disclaimer and references

Titanium tetrachloride or titanium(IV) chloride is the chemical compound with the formula TiCl4.

TiCl4 is an important intermediate in the production of titanium metal and other titanium compounds. It is an unusual example of a liquid metal halide that is very volatile in air, where it forms spectacular opaque clouds of titanium dioxide (TiO2) and hydrogen chloride (HCl). It is sometimes humourously referred to as "tickle".[1]


Properties and structure

TiCl4 is a dense, colourless distillable liquid, although crude samples may be yellow or even red-brown. It is one of the rare transition metal chlorides that is in liquid state at room temperature, VCl4 being another example. This distinctive property arises from the fact that TiCl4 is molecular; that is, each TiCl4 molecule is relatively weakly associated with its neighbours. Most metal chlorides are polymers, where the chloride atoms bridge between the metals. The attraction between the individual TiCl4 molecules is weak, primarily van der Waals forces, and these weak interactions result in low melting and boiling points, similar to those of CCl4.

TiCl4 is tetrahedral, which is consistent with its description as Ti4+ surrounded by four Cl- ligands. Ti4+ has a "closed" electronic shell, with the same number of electrons as the inert gas argon. This configuration leads to highly symmetrical structures, hence the tetrahedral shape of the molecule.

TiCl4 is soluble in toluene and chlorocarbons, as are other non-polar species. Evidence exists that certain arenes form complexes of the type [(C6R6)TiCl3]+. TiCl4 reacts exothermically with donor solvents such as THF to give hexacoordinated adducts.[2] Bulkier ligands (L) give pentacoordinated derivatives TiCl4L.

The main problem with handling TiCl4, aside from its tendency to release corrosive hydrogen chloride, is the formation of titanium oxides and oxychlorides that cement stoppers and syringes.


TiCl4 is produced by the Chloride process, which involves the reduction of titanium oxide ores, typically ilmenite or rutile, with carbon under flowing chlorine at 900 °C. Impurities are removed by distillation to afford pure TiCl4.

2FeTiO3 + 7Cl2 + 6C → 2TiCl4 + 2FeCl3 + 6CO

TiCl4 is inexpensive, thus it is typically purchased for laboratory operations.


Production of titanium metal

TiO2 + 2Cl2 + 2C → TiCl4 + 2CO

Reduction of TiCl4 using magnesium metal produces titanium metal; this is in fact the final step of the Kroll process.

2Mg + TiCl4 → 2MgCl2(l) + Ti

Liquid sodium has sometimes been used instead of magnesium as the reducing agent in the production of titanium metal.

4Na + TiCl4 → 4NaCl + Ti

Production of titanium dioxide

Around 90% of the TiCl4 production is used to make pigment; titanium(IV) oxide (TiO2). Key is the reaction of TiCl4 with water to form hydrochloric acid:

TiCl4 + 2H2O → TiO2 + 4HCl

Or sometimes it is oxidised directly with pure oxygen:

TiCl4 + 2O2 → TiO2 + 2Cl2


In the past titanium tetrachloride has also been used to create naval smokescreens. When sprayed into the air, TiCl4 rapidly reacts with atmospheric moisture:

TiCl4 + 2H2O → TiO2 + 4HCl

The hydrogen chloride immediately absorbs more water to form tiny droplets of hydrochloric acid, which (depending on humidity) may absorb still more water, to produce large droplets that efficiently scatter light. In addition, the intensely white titanium dioxide is also an efficient light scatterer. Due to the corrosiveness of this smoke, however, TiCl4 is no longer used.

Chemical reactions

Organometallic and inorganic chemistry

TiCl4 adopts similar structures to TiBr4 and TiI4; the three compounds share many similarities. TiCl4 and TiBr4 react to give mixed halides TiCl4-xBrx, where x = 0, 1, 2, 3, 4. Magnetic resonance measurements also indicate that halide exchange is also rapid between TiCl4 and VCl4.[3]

TiCl4 is a superb and versatile Lewis acid, as indicated by its tendency to hydrolyze, which implicates the intermediacy of TiCl4(H2O). With THF, TiCl4 forms yellow crystals of TiCl4(THF)2. With Cl- donors, TiCl4 reacts to form sequentially [Ti2Cl9]-, [Ti2Cl10]2-, and [TiCl6]2-.[4] Interestingly, the reaction of chloride ions with TiCl4 depends on the counterion. NBu4Cl reacts with TiCl4 to give the pentacoordinate complex NBu4TiCl5, whereas smaller NEt4 reacts to give (NEt4)2Ti2Cl10. These reactions highlight the influence of electrostatic forces on the structures of compounds with highly ionic bonding.

Much of the extensive organometallic chemistry of titanium starts from TiCl4. Its most important reaction is with sodium cyclopentadienyl to give titanocene dichloride, TiCl2(C5H5)2. This compound is used in organic synthesis (Tebbe's reagent). Arenes, such as C6(CH3)6 reacts to give [Ti(C6(CH3)6)Cl3]+, which is a piano-stool complex.[5] This reaction illustrates the extraordinary Lewis acidity of the TiCl3+ entity, which is derived from TiCl4 using the even stronger chloride-abstracting agent AlCl3.

TiCl4 reacts with four equivalents LiNMe2 to give Ti(NMe2)4, a yellow, benzene-soluble liquid.[6] This molecule is tetrahedral, with planar nitrogen centers.[7]

Reagent in organic synthesis

It is widely used in organic synthesis as a Lewis acid,[8] for example in the Mukaiyama aldol reaction. Key to this application is the tendency of TiCl4 to interact with aldehydes, RCHO, to give adducts such (RCHO)TiCl4OC(H)R. It is also used in the McMurry reaction in conjunction with Zn, LiAlH4, or another reducing agent in order to join two carbonyls in making a carbon-carbon double bond.

Olefin polymerisation

This compound and many of its derivatives are important precursors to Ziegler-Natta catalysts.


Reduction of TiCl4 yields TiCl3. Reduction of TiCl4 with aluminium in THF results in the light-blue THF-adduct TiCl3(THF)3.

Toxicity and safety considerations

Given the tendency of TiCl4 to hydrolyze, the hazards generally arise from the effect of hydrogen chloride. TiCl4 is a strong Lewis acid, exothermically forming adducts with even weak bases such as THF and explosively with water, again releasing HCl.


  1. ^ [1] American Chemistry
  2. ^ L. E. Manzer (1982). "Tetrahydrofuran Complexes of Selected Early Transition Metals". Inorganic Synthesis 21: 135-40.
  3. ^ S. P. Webb, M. S. Gordon (1999). "Intermolecular Self-Interactions of the Titanium Tetrahalides TiX4 (X = F, Cl, Br)". J. Am. Chem. Soc. 121: 2552-2560. doi:10.1021/ja983339i.
  4. ^ C. S. Creaser , J. A. Creighton (1975). "Pentachloro- and Pentabromotitanate(IV) ions". Journal of the Chemical Society, Dalton Transactions: 1402-1405. doi:10.1039/DT9750001402.
  5. ^ F. Calderazzo, I. Ferri, G. Pampaloni, S. Troyanov (1996). "η6-Arene Derivatives of Titanium(IV), Zirconium(IV) and Hafnium(IV)". Journal of Organometallic Chemistry 518: 189-196. doi:10.1016/0022-328X(96)06194-3.
  6. ^ D. C. Bradey, M. Thomas (1960). "Some Dialkylamino-derivatives of Titanium and Zirconium". Journal of the Chemical Society: 3857-3861. doi:10.1039/JR9600003857.
  7. ^ M. E. Davie, T. Foerster, S. Parsons, C. Pulham, D. W. H. Rankin, B. A. Smart (2006). "The Crystal Structure of Tetrakis(dimethylamino)titanium(IV)". Polyhedron 25: 923-929. doi:10.1016/j.poly.2005.10.019.
  8. ^ L.-L. Gundersen, F. Rise, K. Undheim (2004). "Titanium(IV) chloride", Encyclopedia of Reagents for Organic Synthesis, L. Paquette, J. Wiley & Sons. 

General reading

  • Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  • Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Elements, 2nd Edition, Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4. 
This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Titanium_tetrachloride". A list of authors is available in Wikipedia.
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