6518
J. Phys. Chem. 1996, 100, 6518-6523
Oxygen-Bridged Adducts Formed by Chemical Reactions of Solid Aluminum Halides. Experimental Dilemma and ab Initio Calculations of Heterodimer Complexes AlX3-Y (X ) F, Cl; Y ) OH-, H2O) G. Scholz,* R. Sto1 sser, and J. Bartoll Institute of Chemistry and Institute of Physical and Theoretical Chemistry, Humboldt UniVersity of Berlin, Hessische Strasse 1-2, D-10115 Berlin, Federal Republic of Germany ReceiVed: NoVember 2, 1995; In Final Form: January 17, 1996X
Starting from new experimental facts given by ESR, MS, and other physical and chemical methods, the role of bridged intermediates in the formation of catalytic active solids by thermal chemical reactions including the gaseous phase is discussed. Structures of oxygen-bridged adducts of AlF3 and AlCl3 with OH- and H2O have been explored at the HF and MP2 levels using 6-31G*, 6-31+G*, and 6-311G** basis sets. The reaction enthalpies of complex formation are calculated for several temperatures. Both hydroxyl and water complexes were found to be energetically stable complexes that have a real chance to take part in high-temperature heterogeneous reactions like the formation of activated aluminum halide catalysts.
1. Introduction The discovery of the dimeric nature of aluminum and iron trichloride vapors in 1857 has been considered an exception of complex formation for a very long time. Only the development of high-temperature techniques along with mass spectrometric examinations of formed gases since the middle of this century allowed the realization that complex formation in the vapor phase is a general and basic chemical phenomenon. Today, vapor phase complexes are essential in catalysis, chemical synthesis, and high-temperature technical processes. Formally, the formation of 1:1 heterodimer vapor phase complexes can be described by the general equation
MXn(s,1,g) + AYm(g) f MAXnYm(g)
(1)
Such complexes can mainly be formed during evaporation of condensed phases and by solid-gas phase chemical reactions. They are observable above or in coexistence with solids and molten salts and are able to transmit chemical reactions through grain boundaries. The use of “adapted” gaseous intermediates in the formation of activated solids (cf. Figure 1) can in combination with dopings open new ways for the preparation of catalysts (e.g. β-AlF3). It is the aim of this paper, starting from a critical analysis of experimental findings known so far, to extract the basic physical and chemical problems and to give arguments for the existence and nature of the species mentioned in the title. On the basis of quantum chemical ab initio calculations it is the main point to determine the stabilities and equilibrium structures of relevant oxygen-bridged intermediates. Thermodynamic data of the complex formation as well as vibrational frequenices of the complexes are predicted. 2. Experimental Challenge and Dilemma Although there is much indirect evidence for the existence of intermediates in solid state reactions with participation of H2O, there are only a few instances of direct evidence or qualified quantum chemical calculations. A large number of heterodimer halogen-bridged vapor phase molecules containing AlCl3 monomers have been described in X
Abstract published in AdVance ACS Abstracts, March 15, 1996.
0022-3654/96/20100-6518$12.00/0
Figure 1. General scheme for the formation of catalytic active solids. Applied methods: (A) X-ray, DTA, MS, IR, ESR (observation of local structural changes); (B) quantum chemical ab initio calculations; (C) chemical doping during precipitation; simulation of the ESR spectra of doped samples. X′, X: educts and products of the catalytic reaction.
the literature.1 The number of known halogen-bridged heterodimer vapor phase complexes with AlF3 is substantially smaller.2-13 Most of the heterodimers are chemical stable fluoro- or chloroaluminate complexes containing the aluminum halides and further main group or transition metal halides, respectively. Only a few Al-O complexes or aluminum oxide molecules have been successfully investigated experimentally so far. Such experiments suffer from close packed oxidic structures. High melting points of aluminum oxides as well as their low tendency to sublime prevent vaporation of these substances in a feasible amount. The existence of, for example, AlO,14 Al2O, Al3O,14,15 and Al2O216,17 was mainly proven by mass spectrometric, IR, and Raman examinations. In the case of Al2O418 the detection of the reaction of matrix-isolated educts (Al, O2) in inert matrices is the preferred way. Several aluminum oxide halides like AlOF and (AlOF)2 have been described by matrix IR measurements.19 Direct experimental information is available for only two of the vapor phase complexes considered in this paper: the HFAlF3 molecule has been detected by dynamic thermal gas analysis coupled with mass spectrometry;20 the HCl-AlCl3 complex has been evidenced by mass spectrometry too.21,22 Quantum chemical ab initio calculations have shown that all HX-AlX3 complexes (X ) F, Cl) can be described as weak electron pair donor-acceptor adducts with a long and weak intermolecular halogen bridge.23-26 The HF and HCl molecules primarily act as electron donors rather than proton donors in these complexes. The Cl atom in HCl is a weaker Lewis base than the F atom in HF.25,26 © 1996 American Chemical Society
Oxygen-Bridged Adducts AlX3-Y
J. Phys. Chem., Vol. 100, No. 16, 1996 6519
Figure 2. (a) Equilibrium structure (Cs) of HX-AlX3 heterodimers (X ) F, Cl): (cross-hatched circle) Al; (open circles and half-cross-hatched circle) X; (solid circle) H. (b) Structure models for the AlX3-OH- complexes (X ) F, Cl).
TABLE 1: Bond Distances and Anglesa of the Subunits AlCl3, AlF3, OH-, and H2O Calculated with Different Basis Sets molecule/ parameter AlF3 R(AlF) AlCl3 R(AlCl) OHR(OH) H2O R(OH) ∠HOH a
6-31G* HF MP2
6-31+G* HF MP2
6-311G** HF MP2
162.0
164.5
162.7
165.8
162.9
164.9
207.7
207.2
207.7
207.2
207.1
206.6
96.2
98.1
96.2
97.8
95.2
96.8
94.7 105.5
96.9 103.9
94.8 106.5
97.1 105.5
94.1 105.5
95.8 102.5
Distances in picometers, angles in degrees.
TABLE 2: Relative Energies (kJ/mol) of Optimized Structures of AlF3-OH- and AlCl3-OH- at Different Levels of Theorya,b AlF3-OH-
AlCl3-OH-
basis
structure
HF
MP2
HF
MP2
6-31G*
I(Cs) II(Cs) III(C3V) I(Cs) II(Cs) III(C3V) I(Cs) II(Cs) III(C3V)
0.0 (0) 1.3 (1) 420.6 (2) 0.0 (0) 0.4 (1) 404.4 (2) 0.0 (0) 0.6 (1) 423.8 (2)
0.0 1.9 415.8 0.0 0.5 386.1 0.0 1.0 402.3
1.1 (0) 0.0 (1) 461.2 (2) 0.06 (1) 0.0 (0) 438.7 (2) 0.3 (1) 0.0 (0) 456.6 (2)
0.0 0.3 444.4 (2) 0.6 0.0 410.3 0.5 0.0 426.8
6-31+G* 6-311G**
a Number of imaginary frequencies in parentheses. b Total energies (au) of the equilibrium structures. AlF3-OH- (I(Cs)): 6-31G*, -616.006 835 (HF), -616.771 632 (MP2); 6-31+G*, -616.043 829 (HF), -616.839 728 (MP2); 6-311G**, -616.138 007 (HF), -617.028 150 (MP2). AlCl3-OH- (II(Cs)): 6-31+G*, -1696.164 955 (HF), -1696.801 647 (MP2); 6-311G**, -1696.279 809 (HF), -1696.952 085 (MP2).
Direct mass spectroscopic evidence of oxygen-bridged vapor phase complexes such as AlCl3-OH-, AlCl3-H2O or AlF3OH-, and AlF3-H2O was not described in the literature until now. However, there is much empirical chemical knowledge and many spectroscopic results that imply the existence and reaction of such complexes.
Figure 3. Equilibrium structure (Cs symmetry) of the AlF3-H2O and AlCl3-H2O complexes.
From a chemical viewpoint several processes yield a probability for the existence and chemical activity of AlCl3-OH-, AlCl3-H2O or AlF3-OH-, and AlF3-H2O complexes. Hygroscopic properties of solid AlCl3 (or FeCl3) should favor the formation of AlCl3-H2O (FeCl3-H2O) or the corresponding hydroxyl complexes during heating and melting processes even in the presence of only traces of water and at abient temperatures. (Rustad et al.27 formulated spectrophotometric evidence of iron(III) chloride hydrate vapor molecules FeCl3-H2O.) The complex nature of thermal phase transitions in aluminum fluoride hydrates28,29 gave strong evidence for the formation of AlF3-OH- or AlF3-H2O complexes. (The existence of HFAlF3 and 2HF-AlF3 was directly proven during such processes.20,28) Thermally induced processes such as water loss, hydrolysis occurring after the main water content was given up, and the formation of the highly reactive species HF accompanied by a complex chemical and physical reorganization of the solid cause the appearance of a spontaneous catalytic active β-AlF3 phase above 450 °C.28-30 Menz et al.28 assumed that the intermediate formation of solid and gaseous aluminum fluoride oxides plays a crucial role for the reorganization processes taking place in the solid. F/OH- exchange processes could be proven during dehydration processes.28 Most of the questions concerning the formation of the catalytic active phase during thermal decomposition of the educts, the role of vapor phase complexes for that process, and the knowledge of the nature and structure of active sites are still open. A consistent description of the complex thermal reaction of aluminum fluoride hydrates under different experimental conditions does
6520 J. Phys. Chem., Vol. 100, No. 16, 1996
Scholz et al.
TABLE 3: Optimized Geometries for the Cs Structures I of AlF3-OH- and II of AlCl3-OH- a molecule/ basis set
R(AlO) HF/MP2 R(OH) HF/MP2 R(AlX′) HF/MP2 R(AlX) HF/MP2 R(HX)b HF/MP2 ∠AlOH HF/MP2 ∠OAlX HF/MP2
AlF3-OH6-31G* G-31+G* 6-311G** AlCl3-OH6-31G* 6-31+G* 6-311G** a
175.5/177.9 175.5/178.3 174.4/176.7
94.6/96.8 94.5/96.9 93.7/95.6
168.0/170.2 169.2/172.3 169.4/171.2
168.8/171.0 170.2/173.3 170.1/171.8
268.2/263.3 285.4/278.2 277.0/266.0
110.8/106.4 115.0/112.4 115.1/107.8
106.8/106.4 107.6/107.7 107.8/107.1
172.7/175.4 173.2/176.3 172.1/174.5
94.5/96.8 94.5/97.0 93.7/95.7
219.5/218.0 219.3/217.6 218.9/217.3
217.4/216.0 217.2/215.6 216.8/215.4
342.4/340.8 344.7/343.6 345.5/337.6
116.7/113.1 118.6/115.8 120.0/112.4
109.6/108.8 109.5/108.2 109.6/109.3
Distances in picometers, angles in degrees. b Hydrogen bridge to the nearest halogen atom.
TABLE 4: Energies of Complex Formation AlX3 + Y S AlX3-Y for AlX3-OH- and AlX3-H2Oa AlF3-OH-b ∆E ∆Ec ∆ AlCl3-OH-c ∆E ∆Ec ∆ AlF3-H2O ∆E ∆Ec ∆ AlCl3-H2O ∆E ∆Ec ∆
6-31G*
6-31+G*
6-311G**
-603.3 -522.6 80.7
-517.1 -509.0 8.1
-615.1 -535.1 80.0
-665.5 -576.7 88.8
-552.7 -541.1 11.6
-660.5 -572.9 87.6
-143.4 -122.6 20.8
-133.7 -124.4 9.3
-148.4 -130.0 18.4
-129.4 -112.4 17.0
-116.9 -105.5 11.4
-129.4 -112.9 16.5
evaluation of electron correlation effects has been performed by MP2(full) optimizations of all geometries as well. Three basis sets have been employed at both the HF and the MP2 level: (i) 6-31G*, split valence basis set plus polarization functions on Al, F, O, Cl;34,35 (ii) 6-31+G*, split valence plus polarization functions and diffuse sp functions on Al, F, O, Cl;34-36 (iii) 6-311G**, split valence basis set plus polarization functions on all atoms.37,38 The thermodynamic functions of complex formation were calculated within the standard rigid-rotor harmonic-oscillator ideal-gas approximation.34 The calculations have been performed on IBM RS6000 and HP9000/735 workstations using the Gaussian 9239 program. 4. Results and Discussion
a
HF level, ∆E without and ∆Ec with corrections for the basis set superposition error in kJ/mol; basis set superposition error ∆ in kJ/ mol. b Values for the equilibrium structure I. c Values for the equilibrium structure II.
not exist in the literature. To answer those questions, we mainly used two ways: (i) a theoretical one by performing quantum chemical ab initio calculations of selected vapor phase species like bridged complexes; (ii) an experimental one using hightemperature ESR in situ spectroscopy. Paramagnetic probes (Cu2+, Mn2+, Fe3+, Cr3+) are incorporated into the fluoride solids,31,32 and the local and overall structural changes are observed in dependence on temperature and in combination with diffraction and thermal methods. 3. Methods Although the application of the density functional theory in molecular quantum mechanics plays an important role, the good quality of HF and MP2 calculations combined with extended basis sets still favor their application for molecules, molecular clusters, and complexes.33 Geometries of different structure models of AlX3-OH- and AlX3-H2O complexes have been optimized at the HartreeFock level using ab initio molecular orbital theory.34 The
Bond distances and bond angles of the isolated species considered in this paper are summarized in Table 1. For the determination of the most stable structure of HF-AlF3 five structures (Figures 1 and 224) have been optimized. It could be shown that the only local minimum on the potential energy surface of HF-AlF3 and all similar HX-AlX3 molecules is the Cs structure shown in Figure 2a.24,26 This structure consists of a slightly pyramidal AlX3 molecule and a nearly undistorted HX molecule, which is bonded to AlX3 through both the halogen and the hydrogen in a cyclic arrangement. AlX3-OH-. The three most plausible structure models used for the calculations on AlX3-OH- complexes are illustrated in Figure 2b. Relative energies of all three structures and the number of negative frequencies are listed in Table 2. The results obtained using all basis sets mentioned above indicate that structures II and III are saddle points, whereas structure I is the minimum structure of AlF3-OH-. This result is parallel to the results on halogen-bridged dimers mentioned above. The energetic difference between the two Cs structures I and II of HF-AlF3 covers the range 4.1-12.9 kJ/mol, depending on basis sets and methods used.24 This difference ∆EI-II between the local minimum and the first-order saddle point, which corresponds to the rotation barrier of the HX subunit around the intermolecular bridge, was found to be extremely diminished
TABLE 5: Optimized Geometries of the Equilibrium Structures of AlF3-H2O and AlCl3-H2O Complexesa molecule/ basis set AlF3-H2O 6-31G* 6-31+G* 6-311G** AlCl3/H2O 6-31G* 6-31+G* 6-311G** a
R(AlO) HF/MP2
R(OH) HF/MP2
R(AlX′) HF/MP2
R(AlX) HF/MP2
R(HX′) HF/MP2
∠AlOH HF/MP2
∠AlOX HF/MP2
∠HOH HF/MP2
196.0/198.0 195.2/197.4 193.7/196.5
95.4/97.7 95.4/97.8 94.6/96.4
164.4/166.8 165.3/168.4 165.4/167.3
163.5/165.8 164.4/167.3 164.7/166.4
335.0/330.6 338.6/339.2 342.1/336.9
113.4/108.3 115.1/112.3 118.2/113.2
106.1/107.0 105.5/105.4 105.4/105.7
109.0/107.4 109.0/107.9 110.2/107.5
196.4/198.4 196.5/197.1 194.8/196.0
95.5/97.8 95.4/97.5 94.6/96.2
212.0/211.2 211.9/210.6 211.0/210.0
210.5/209.5 210.5/210.2 210.7/209.7
382.2/375.6b 382.9/376.1 395.7/396.5
117.8/113.1 118.2/124.6 124.3/124.8
104.6/104.8 104.4/103.2 103.6/103.7
109.6/108.0 109.5/110.8 111.3/110.3
Distances in picometers, angles in degrees. b The hydrogen bonds to the two chlorines are slightly different.
Oxygen-Bridged Adducts AlX3-Y
J. Phys. Chem., Vol. 100, No. 16, 1996 6521
TABLE 6: Enthalpies of Complex Formation ∆HT (kJ/mol) and Equilibrium Constants K (atm-1) of AlF3-H2O and AlCl3-H2O for the Reaction AlX3 + H2O S AlX3-H2Oa AlF3-H2O
AlCl3-H2O
T (K)
∆HT
K
∆HT
K
273 498 598 648 698 748 1148
-117.7 -116.7 -115.8 -115.3 -114.8 -114.3 -109.9
3.9 × 1013 2.8 ×103 2.6 × 10 4.3 0.9 0.2 4.6 × 10-4
-98.0 -95.8 -94.5 -93.8 -93.1 -92.4 -86.3
1.4 × 1012 5.6 × 103 1.2 × 102 2.8 × 10 8.1 2.8 0.02
a Calculations with the HF/6-31+G* equilibrium geometries, vibrational frequencies, and BSSE-corrected dissociation energies of the complexes. Zero-point vibrational energies: AlF3-H2O, 11.2 kJ/mol; AlCl3-H2O, 10.2 kJ/mol.
in the case of oxygen-bridged complexes. It ranges from 0.4 to 1.9 kJ/mol for AlF3-OH- and becomes more pronounced with the inclusion of correlation effects. Therefore, the assumption of a nearly free rotation of the OH group around the fixed Al-O axis should be allowed. AlCl3-OH- was found to fit structure II with 6-31+G* and 6-311G** basis sets at the HF and MP2 levels. However, the highest energetic difference between structures I and II is 0.6 kJ/mol at the MP2/6-31+G* level, which implies a more flat potential energy surface for the rotation of the OH group of this complex. Obviously, the 6-31G* basis (lacking a polarization function on H and diffuse sp functions on Al, Cl) is not suited for the description of the chlorine-containing complex, whereas good results could be obtained for the individual AlCl3 molecule using this basis set.40 Table 3 contains the optimized geometries of the equilibrium structures of both complexes obtained at the HF and MP2 levels. AlF3-OH- and AlCl3-OH- complexes consist of only slightly distorted OH- groups with shortened OH distances and pyramidal AlF3 (AlCl3) moieties with AlF (AlCl) bonds remarkably larger than in free AlF3 and AlCl3 molecules, respectively (cf. Table 1). MP2 calculations led to an increase of all bond lengths of AlF3-OH-, even of the intermolecular bond. This tendency is opposite to that of weak intermolecular halogen-bridged dimers, where an increase of the bond lengths has been observed only within the subunits.24,25 However, in all cases of halogen-bridged dimers the inclusion of electron correlation effects resulted in shorter intermolecular distances.24,25 A further reversed effect can be observed in AlCl3OH- with diminished Al-Cl bond lengths and an enlarged
intermolecular Al-O bond at the MP2 level (Table 3). In agreement with halogen-bridged dimers considered so far electron correlation has a diminishing influence on the intermolecular angle {AlOH} and on the hydrogen-halogen bridge as well. The intermolecular Al-O distance in AlF3-OH- is only about 5 pm longer than the AlF bond lengths. This results in an approximately tetrahedral structure of the {AlF3-O}2unit. The calculated intermolecular Al-O distances are remarkably smaller than corresponding intermolecular Al-F or Al-Cl distances in vapor phase complexes.23-26 They are in good agreement with Al-O distances in neutral Al-H2O radical complexes calculated using density functional theory.41 Due to collective interactions in solids, the same tendency can be observed for Al-O and Al-F (Al-Cl) bond lengths in solid aluminum oxides and halides, respectively. The {AlOH} angles are widened compared with {AlFH} or {AlClH} angles.24-26 Whereas the AlX3 subunits in the halogen-bridged dimers remain nearly planar, they distinctly show a pyramidal structure in the oxygen-bridged dimers. These differences in the structures of halogen- and oxygen-bridged complexes have interesting consequences for their complexation energies, summarized in Table 4, as well as for their overall chemical properties. The energies of complex formation corrected for the basis set superposition errors of HX-AlX3 (X ) F, Cl) complexes lie between -13.0 kJ/mol (HCl-AlCl3) and -67.0 kJ/mol (HFAlF3).24,25 With respect to their binding energies the fluorinebridged complexes represent typical examples of species that are characterized by interactions covering the range between chemical and intermolecular nature. In comparison, chlorinebridged complexes are typical representatives of van der Waals complexes. AlF3-OH- and AlCl3-OH- complexes, however, are energetically stable complexes with corrected interaction energies below -500 kJ/mol. Even in the case of addition of a further proton (AlF3-H2O, AlCl3-H2O), the complexes are much more stable than analogous halogen-bridged complexes. AlF3-H2O is about twice and AlCl3-H2O is about 1.6 times more stable than the most stable fluorine-bridged complex, HF-AlF3. Obviously, the inclusion of diffuse sp functions has a pronounced influence on the interaction energies of the charged hydroxyl complexes. The superposition errors could be reduced to ∼8-11 kJ/mol using 6-31+G* basis set compared to ∼8090 kJ/mol for the basis sets without diffuse functions (cf. Table 4).
TABLE 7: Vibrational Frequencies (cm-1) of the Equilibrium Structures of AlF3-OH-, AlCl3-OH-, AlF3-H2O, and AlCl3-H2O (6-31+G* Basis Set, HF Level) AlF3-OH-
AlCl3-OH-
mode
νi
intens.a
intermol. intermol. δFAlF δFAlF δFAlF umbr. intermol. intermol. (νAlO) intermol. νAlF νAlF νAlF νOH
190 212 313 313 320 525 679 806 851 853 883 3950
0.3 0.4 11.4 11.3 15.5 34.2 0.4 14.9 100.0 59.9 49.3 7.8
a
Relative IR intensities.
AlF3-H2O
AlCl3-H2O
mode
νi
intens.a
mode
νi
intens.a
mode
νi
intens.a
intermol. δClAlCl δClAlCl δClAlCl umbr. intermol. intermol. νAlCl νAlCl νAlCl intermol. intermol. (νAlO) νOH
18 136 155 198 222 260 378 480 491 748 810 4155
18.0 0.2 3.2 4.1 5.7 20.2 6.2 95.1 67.6 100.0 34.7 13.8
intermol. intermol. intermol. δFAlF δFAlF δFAlF umbr. intermol. (νAlO) intermol. νAlF intermol. νAlF νAlF δHOH νOH νOH
115 162 192 253 264 300 472 510 710 731 944 949 1806 4003 4102
1.6 5.0 6.0 9.2 14.6 38.0 26.0 100.0 10.1 3.0 79.2 76.0 48.0 38.6 78.7
intermol. δClAlCl intermol. δClAlCl intermol. δClAlCl umbr. νAlCl intermol. intermol. (νAlO) νAlCl νAlCl intermol. δHOH νOH νOH
60 127 160 179 181 206 390 415 451 562 603 757 1806 3991 4093
2.2 3.2 8.7 3.0 0.2 16.6 5.3 72.8 27.6 39.2 100.0 44.0 58.0 51.0 81.0
6522 J. Phys. Chem., Vol. 100, No. 16, 1996 AlX3-H2O. It is an interesting fact that according to the hitherto existing knowledge on water and HX (X ) F, Cl)23-25,42-45 containing intermolecular complexes with AlX3 molecules, only one structure (given in Figure 3) was expected and obtained as the equilibrium structure for this type of complex. In principle this structure is similar to those obtained for AlF3-OH- (Figure 2b, structure I(Cs)) with the second proton situated in the neighbored OAlX-plane. This means that two of the three fluorines are perfectly eclipsed by the hydrogens. In AlCl3-H2O the {HOH} angle is slightly widened, with the consequence of a non-perfectly eclipsed structure. The optimized bond lengths and angles for the water complexes are given in Table 5. They confirm very recently published results of Ball45 obtained with one basis set (6-31G(d,p)) and demonstrate the dependence on several extended basis sets used in this paper. The inclusion of electron correlation effects (MP2) has the same influence as described for the OHcomplexes. The intermolecular distances are ∼20 pm longer than in the charged hydroxyl complexes. They are in the same range as for the HF-AlF3 complex. These geometries are properly reflected by the complexation energies (Table 4) discussed above. In agreement with Ball45 the AlF3-H2O complex was found to be more stable than the AlCl3-H2O complex. Furthermore, the complex formation between two neutral monomers has a less pronounced influence on the structure of the monomeric subunits (for comparison Tables 3 and 5) with respect to the charged hydroxyl complexes. The calculation of the thermodynamic functions of complex formation was performed with the basis set giving the lowest superposition error. Therefore, the enthalpies of complex formation, ∆HT, of all complexes were calculated with the corrected HF/6-31+G* dissociation energies and the HF/631+G* equilibrium geometries. For the hydroxyl complexes ∆HT was obtained as AlF3-OH-, -501.5 kJ/mol (498 K), -500.8 kJ/mol (698 K); and AlCl3-OH-, -533.2 kJ/mol (498 K), -532.2 kJ/mol (698 K). The values for the water complexes were summarized together with their equilibrium constants in Table 6. These results clearly indicate that in the temperature range of reconstructive phase transitions of aluminum fluoride hydrates31,32 (373 K e T e 773 K) even the water complex with AlF3 is more stable than the detected HF-AlF3 complex24 (HF-AlF3: ∆H(498 K) ) -60.4 kJ/mol, ∆H(698 K) ) -58.4 kJ/mol). The vibrational frequencies of the equilibrium structures of all hydroxyl and water complexes are collected in Table 7. The normal vibrations of all complexes could be assigned with the help of the program Viewmol.47 The vibrations of the hydroxyl complexes are strongly coupled, and the existence of several higher intermolecular modes underline their energetic stability. 5. Conclusions AlF3-OH-, AlCl3-OH-, AlF3-H2O and AlCl3-H2O complexes are found to be energetically stable complexes. The complexation energies are about 1.6-8 times more stable than the most stable fluorine bridged complex, HF-AlF3. The latter complex, despite being less stable, was detected unambigously by mass spectroscopy as a result of hydrolysis processes accompanying the thermal decomposition of water containing aluminum fluorides. The assumption of the existence of AlF3H2O or AlF3-OH- seems to be realistic, although direct experimental evidence is not available as yet. Among other processes, their formation could play a decisive role for thermally induced reorganization processes in aluminum fluoride hydrates. Furthermore, the formation of AlCl3-H2O or analogous FeCl3-H2O adducts has to be taken into account during heating
Scholz et al. and melting processes of AlCl3, FeCl3, or mixtures of them if water traces cannot be excluded. Difficulties for their detection arise from the composition of the vapor phase, consisting of Al2Cl6, Fe2Cl6, FeAlCl6, FeCl3, and AlCl3, in dependence on the temperature and the similar complex stability of the dimers.46 Finally, the existence of adducts of aluminum(III) fluoride and aluminum(III) chloride with water can be predicted as a result of these calculations, and it is still a challenge to the experimentalists to prove their existence. Acknowledgment. The Deutsche Forschungsgemeinschaft is kindly acknowledged for financial support. References and Notes (1) (a) Scha¨fer, H. Angew. Chem. 1976, 88, 775. (b) Bock, Ch. W.; Trachtman, M.; Mains, G. J. J. Phys. Chem. 1993, 97, 2546. (2) Kolosov, E. N.; Sholtz, V. B.; Sidorov, L. N. Vestn. MoskoVsk. UniV. 1972, 1, 49. (3) Vajda, E.; Hargittai, I., Tremmel, J. Inorg. Chim. Acta 1977, 25, L143. (4) Cyvin, S. J.; Cyvin, B. N.; Snelson, J. J. Phys. Chem. 1971, 75, 2609. (5) Huglen, R.; Cyvin, S. J.; Oye, H. A. Z. Naturforsch. 1979, 349, 1118. (6) Sholtz, V. B.; Sidorov, L. N. Vestn. MoskoVsk. UniV. 1972, 4, 371. (7) Snelson, A.; Cyvin, B. N.; Cyvin, S. J. J. Mol. Struct. 1975, 24, 165. (8) Spiridonov, V. P.; Erokhin, E. V. Russ. J. Inorg. Chem. 1969, 14, 332. (9) Sidorov, L. N.; Nikitin, M. I.; Shokan, E. V.; Sorokin, I. D. AdV. Mass. Spectrosc. 1980, 8A, 462. (10) Sorokin, I. D.; Sidorov, L. N.; Nikitin, M. I.; Shokan, E. V. Int. J. Mass Spectrom. Ion Phys. 1981, 41, 45. (11) Morozov, I. V.; Korenev, Yu. M. Zh. Neorg. Khim. 1993, 38, 659. (12) Nikitin, M. I.; Sorokin, I. D.; Shokan, E. V.; Sirodov, L. N. Zh. Fiz. Khim. 1980, 54, 1337. (13) Menz, D.-H.; Scholz, G.; Becker, D.; Binnewies, M. Z. Anorg. Allg. Chem. 1994, 620, 1976. (14) JANAF, Thermochemical Tables; Stull, D. R., Prophet, H., Eds.; U.S. National Bureau of Standards: Washington, DC, 1971. (15) Cox, D. M.; Trevor, D. J.; Whetten, R. L.; Rohlfing, E. A.; Kaldor, A. J. J. Chem. Phys. 1986, 84, 4651. (16) Zaitsevskii, A. V.; Chertilin, G. V.; Serebrennikov, L. V.; Stepanov, N. F. J. Mol. Struct. (THEOCHEM) 1993, 280, 291. (17) Nemukhin, A. V.; Weinhold, F. J. Chem. Phys. 1992, 97, 3420. (18) Nemukhin, A. V. J. Mol. Struct. (THEOCHEM) 1994, 315, 225. (19) Ahlrichs, R.; Zhengyan, L.; Schno¨ckel, H. Z. Anorg. Allg. Chem. 1984, 519, 155. (20) Menz, D.-H.; Kolditz, L.; Heide, K.; Schmidt, C.; Kunert, Ch.; Mensing, Ch.; Schnering, H. G. v.; Ho¨nle, W. Z. Anorg. Allg. Chem. 1987, 551, 231. (21) Rabeneck, H.; Scha¨fer, H. Z. Anorg. Allg. Chem. 1973, 395, 69. (22) Imanaka, T. Bull. Chem. Soc. Jpn. 1967, 40, 2182. (23) Scholz, G.; Sauer, J.; Menz, D.-H. Chem. Phys. Lett. 1989, 156, 125. (24) Curtiss, L. A.; Scholz, G. Chem. Phys. Lett. 1993, 205, 550. (25) Scholz, G. J. Mol. Struct. (THEOCHEM) 1994, 309, 227. (26) Wilson, M.; Coolidge, M. B.; Mains, G. J. J. Phys. Chem. 1992, 96, 4851. (27) Rustad, D. S.; Gregory, N. W. Inorg. Chem. 1988, 27, 2840. (28) Menz, D.-H.; Mensing, Ch.; Ho¨nle, W.; Schnering, H. G. v. Z. Anorg. Allg. Chem. 1992, 611, 107. (29) Menz, D.-H.; Zacharias, A.; Kolditz, L. J. Therm. Anal. 1988, 33, 811. (30) Kemnitz, E.; Hansen, G.; Hess, A.; Kohne, A. J. Mol. Catal. 1992, 77, 193. (31) Sto¨sser, R.; Scholz, G.; Pa¨ch, M. J. Solid State Chem. 1995, 116, 249. (32) Scholz, G.; Sto¨sser, R.; Sebastian, S.; Kemnitz, E.; Bartoll, J. J. Phys. Chem. Solids, submitted. (33) Molecular Quantum Mechanics: Methods and Applications, International Conference in memory of S. F. Boys and in honor of I. ShaVitt, Cambridge, Sept 3-7, 1995, Book of Abstracts. (34) Hehre, W. J.; Radom, L.; Schleyer, P. v. R.; Pople, J. A. Ab Initio Molecular Orbital Theory; Wiley: New York, 1986. (35) Hariharan, P. C.; Pople, J. A. Theor. Chim. Acta 1973, 28, 213. (36) Clark, T.; Chandrashekar, J., Spitznagel, G. W.; Schleyer, P. v. R. J. Comput. Chem. 1983, 4, 294. (37) Krishnan, R.; Binkley, J. S.; Seeger, R.; Pople, J. A. J. Chem. Phys. 1980, 72, 650.
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