Mixed-Valent Titanium Complex

Aqueous Chemistry of Titanium(ii) Species

Ulrich K�lle* and Philipp K�lle

Reports on titanium in the oxidation state + ii are confined tosolid-state compounds, such as the nonstoichiometric oxideTiOx, the hydride TiH2 as well as halides TiF2 and TiCl2 and anumber of coordination compounds [TiCl2Ln]. Although acalculated redox potential of E8(TiIII/TiII)=�0.37 V versusthe normal hydrogen electrode (NHE) appears in textbooks[1]

and tables,[2] which suggests the existence of the TiIIaq ion, theaqueous chemistry of TiII is almost unknown.[1,3]

Titanium metal, though quite electropositive (E8(Ti3+/Ti8=�1.2 V vs. NHE), does not dissolve in acids, exceptaqueous HF.[4] In the course of corrosion studies of titanium inHF Straumanis et al. observed a greenish-yellow complexwhich was assumed to contain the [TiF6]

3� ion.[5] Theoxidation state + iii of the Ti species formed was verifiedby redox titration and by monitoring the amount of hydrogenevolved on titanium metal dissolution.

We found that titanium metal (0.5 mm wire, > 99%)dissolves readily with vigorous hydrogen evolution either inaqueous HF or in aqueous acids, such as HBF4 (5%),[6]

CF3SO3H, or H2SO4 (2m) if these contained HF. Whereasthe color of the resulting solution in HF resembles thatdescribed by Straumanis et al.,[7] a green solution is obtainedin the stronger acids. The absorption spectrum in acid solution(Figure 1) which has two bands at 430 and 650 nm, is not inaccord with a TiIII complex.[8] Instead it is reminiscent of thespectrum of TiII doped into NaCl[9] and closely resembles thatof the V3þ

aq ion (� 400 and 600 nm assigned to 3T1g!3T2g and3T1g!3T2g(P) transitions[10]), which, like the TiII species, has a

Figure 1. Absorption spectrum of Ti2þaq at pH 0.5.

[*] Prof. Dr. U. K�lle, Dipl. Chem. P. K�lleInstitute for Inorganic ChemistryRWTH Aachen52056 Aachen (Germany)Fax: (+49)241-80-92288E-mail: u.koelle@ac.rwth-aachen.de

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4540 � 2003 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim DOI: 10.1002/anie.200351280 Angew. Chem. Int. Ed. 2003, 42, 4540 –4542

d2 configuration.[11] The green solution is moderately airsensitive and slowly turns brownish red (TiIII) and finally tocolorless (TiIV) on exposure to air.

Concentrating the solution to 0 0.2m causes formation ofa gray precipitate. The solid was collected and could beisolated after complete evaporation of water, followed bytreatment with diethyl ether to remove excess HBF4, andwashing with acetone. Analysis of this material is in accordwith the formula TiF3(H2O)2.

[12]

The structure of this compound was solved from X-raypowder data.[13] The Rietveld refinement converged withgood to excellent R values, the observed and calculatedprofile plot is shown in Figure 2. However, standard devia-tions of positional parameters,[13] and therefore of bondlengths are unusually large. This effect might be due tocrystallographic stacking faults of the chain structure, aproblem that cannot be recognized nor resolved frompowder data.

The structure is built up from chains of corner-sharing,trans-bridged, TiX6 octahedra of crystal chemical formulation[TiF4F2/2][Ti(H2O)4F2/2] stacked along the crystallographic caxis (Figure 3). The compound is isostructural with Cu(H2O)4-

TiF6, described by Fischer et al.[14] . In accord with solutionbehavior (see below) the oxidation states are assigned as TiIV

for [TiF4F2/2] and TiII for [Ti(H2O)4F2/2]. Since fluorine andoxygen are almost indistinguishable in terms of their X-rayscattering factors, labels have been assigned according to thetwo sets of Ti�X separations (X=F, O) found for the two Ticenters. The long distance of 2.07(2) > was assigned to theTiII�OH2 bond. This TiII center also has the longer bond to thebridging fluoride, 2.00(2) >. The TiIV center has shorter bondsto terminal, 1.94(1) and bridging 1.97(2) > fluorides. Thechains are interconnected by hydrogen bonds, which issuggested by the shortest interchain contacts between watermolecules of 2.65 >. Adjacent chains are offset along the caxis by 2.62 >, which results in an interlocked packing mode.

The X-band powder EPR spectrum of this material atroom temperature shows two closely spaced lines with thesame phase at g= 1.9469 and 1.9485 with linewidths 120 G

and 24 G, respectively. The spectrum is indicative of a cubicS= 1 system with a very small value of the zero-field splittingparameter D (D< 0.003 cm�1).

The solid slowly dissolves in neutral water to give a clearyellow green solution which on acidification (with HBF4, forexample) turns into bright green and shows the sameabsorption spectrum as that obtained on dissolving titaniumdirectly. We thus explain the dissolution of titanium metal inHF or acids containing HF according to Equation (1).

2 Ti þ 6HF ! ½TiF6�2� þ Ti2þ þ 3H2 ð1Þ

Owing to stoichiometric formation of TiII and TiIV theaverage oxidation state found for titanium in the ensuingsolution, as well as the amount of hydrogen released, is thesame as if TiIII would have been formed, which explains theconclusions drawn by Straumanis et al.[7] Equation (1) furtherimplies disproportionation of TiIII in the presence of excessfluoride according to Equation (2).

2 Ti3þ þ 6F� ! ½TiF6�2� þ Ti2þ ð2Þ

We verified that TiCl3, dissolved in aqueous HF, directlygave the green yellow solution of TiII/TiIV which correspondsto Equation (2). A value of EO([TiF6]

2�/Ti8)��0.8 V[4] fromcorrosion measurements of titanium in HF along with thecalculated value of EO(Ti3+/Ti8)=�1.2 V leads to a verypositive value for EO([TiF6]

2�/Ti3+) and thus to a positivepotential for the disproportionation in Equation (2).

Figure 2. Upper trace: Observed verses calculated profile intensity ofTi(TiF6)(H2O)4 as obtained from Rietveld refinement, lower trace: dif-ference curve, position of the Bragg reflections.

Figure 3. Structure of Ti(TiF6)(H2O)4 viewed along the b axis. TiIVF6octahedra in red, TiIIF2(H2O)4 octahedra in blue.

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Excess F� as well as [TiF6]2� can be precipitated from the

solution with Ba2+ or Ca2+ ions (as BF4� or CF3SO3

� salts)whereby the absorption spectrum of the solution does notchange. Since this spectrum is not influenced by the anion itmust arise from the Ti2+ aqua ion, most probably[Ti(H2O)6]

2+. The formation of an insoluble hydroxide atpH� 7 is typical for a first-row + ii transition-metal ion.Further characterization of the novel aqua ion and itscomplexing behavior is in progress in our laboratory.

Received: February 26, 2003Revised: June 24, 2003 [Z51280]

.Keywords: coordination chemistry · fluorides ·structure elucidation · titanium · UV/Vis spectroscopy

[1] D. T. Richens, The Chemistry of Aqua Ions, Wiley, New York,1997, p. 208.

[2] M. Pourbaix, Atlas d'Equlibres Electrochimiques, Gauthier-Villars, Paris, 1963, p. 214.

[3] F. A. Cotton, G. Wilkinson, Advanced Inorganic Chemistry,Wiley, New York, 5th ed, 1988, p. 651.

[4] Encyclopedia of Electrochemistry of the Elements, Vol. V (Ed.:A. J. Bard), Marcel Dekker, New York, 1976, p. 307, andreferences therein.

[5] M. E. Straumanis, C. H. Cheng, A. W. Schlechten, J. Electro-chem. Soc. 1956, 103, 439; M. E. Straumanis, C. H. Cheng, A. W.Schlechten, J. Electrochem. Soc. 1956, 103, 440.

[6] The HF content of the HBF4 is responsible for the dissolution ofthe titanium.

[7] Absorption spectra of solutions containing excess fluoride arevery pH dependent, the l= 430 nm band increases in intensityand shifts to slightly longer wavelengths at pH> 3. Since thespectrum of the aqua complex does not change up to pH 6.5, theyellowish species obtained at higher pH and observed byStraumanis must be a titanium(ii) fluoro complex.

[8] A d1 complex shows only one band in the visible region, themaximum Jahn–Teller splitting for this band in the case of TiIII is� 60 nm.

[9] TiII, 0.5 atom% in NaCl shows bands at l= 500 and 800 nm inthe UV/Vis spectra; a) W. E. Smith, J. Chem. Soc. Chem.Commun. 1972, 1121; b) D. H. Brown, A. Hunter, W. E. Smith,J. Chem. Soc. Dalton Trans. 1979, 79.

[10] a) ref. [1], p. 232; b) A. B. P. Lever, Inorganic Electronic Spec-troscopy, Elsevier, Amsterdam, 1984, pp. 399 – 401.

[11] From the bands at l= 430 and 650 nm values of B= 432 cm�1

and Doct= 11665 cm�1 are calculated, compare V3þaq B= 665 cm�1

and Doct = 18620 cm�1.[12] (%) Calcd: Ti 34.0, F 40.4; found Ti (photometrically H2O2) 35.7,

F 40.0.[13] STOE Debeye Scherrer diffractometer, 0.3 mm capillary, CuKa

radiation, 15 2V 1008. Structure solution with EXPO,[16]

Rietveld refinement with FULLPROF[15] (pseudo-Voigt profilefunction, 5 profile parameters, 1 parameter for preferredorientation, 2 asymmetry parameters, 1 overall scale parameter,4 parameters for lattice constants, 8 positional parameters, 5isotropic displacement parameters). Rp= 0.028, Rwp= 0.037,RBragg= 0.022, GooF 0.97 for 226 Bragg reflections and 27refined parameters. Space group: C2/m ; lattice constants: a=5.7075(2), b= 10.2193(3), c= 7.8915(2) >, b= 117.421(1)8 ; frac-tional coordinates and isotropic displacement parameters: TiII 00 0, Biso 1.6(2) >2; TiIV 0 0 1=2, Biso 1.1(2) >2; F(1) 0.998(5) 00.746(3), Biso 2.4(1) >2; F(2) 0.732(2) 0.861(1) 0.411(1), Biso

1.6(2) >2; O 0.708(2) 0.858(1) 0.900(2), Biso 3.7(4) >2. Inter-

atomic distances [>]: TiII-F(1) 2.00(2), TiII-O 2.07(1), TiIV-F(1)1.94(2), TiIV-F(2) 1.97(1).

[14] J. Fischer, G. Keib, R. Weiss, Acta Crystallogr. 1967, 23, 338.[15] J. Rodriguez-Cavajal FULLPROF 2000, Version 1.9b Labora-

toire LKon Brillouin 2000.[16] C. Giacovazzo, D. Siliqui, B. Carrozzini, A. Guagliardi, A.

Moliterni, J. Appl. Crystallogr. 1999, 32, 115.

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