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DALTON J. Chem. Soc., Dalton Trans., 1997, Pages 1953–1956 1953 Reactions of ·-lithiated phosphinimines with PhCN; the crystal structure of [K{N(H)C(Ph)C(H)P(Ph)2] NSiMe3}(tmen)]2 (tmen = Me2NCH2CH2NMe2) Peter B. Hitchcock, Michael F. Lappert * and Zhong-Xia Wang The Chemistry Laboratory, University of Sussex, Brighton BN1 9QJ, UK Treatment of the a-lithiated phosphinimine Li{CH(R9)P(R)2]] NSiMe3} with benzonitrile yielded (via a trimethylsilyl or hydrogen 1,3 C æÆ N shift) the trimethylsilyliminophosphoranylenamidolithium complex Li{N(R9)C(Ph)C(H)P(R)2]] NSiMe3} (R = Me, R9 = SiMe3 1; R = Ph, R9 = SiMe3 2; or R = Ph, R9 = H 3).Complex 2 was transformed into the corresponding potassium complex 4 by an exchange reaction with KOBut. Crystallisation of 4 from hexane in the presence of Me2NCH2CH2NMe2 (tmen) gave the tmen adduct 5 and a trace of the partially hydrolysed product [K{N(H)C(Ph)C(H)P(Ph)2]] NSiMe3}(tmen)]2 6, which was characterised by a single crystal X-ray diffraction study as a dinuclear complex with each of the two potassium atoms in a different co-ordination environment. Complex 1 or 2 was hydrolysed to form the neutral iminophosphoranylenamine N(SiMe3)C(Ph)C(H)P(R)2N(SiMe3)H 7 or 8, which showed (1H NMR spectroscopy) the presence of hydrogen bonds.Reactions of a-lithiated phosphinimines with several electrophiles have been explored. Thus, when an a-lithiated phosphinimine was treated with Me3SiCl, Br2 or certain organic electrophiles, a carbon-centred reaction was invariably observed.This was ascribed to the difference in nucleophilicity between the two ylidic positions.1 Such reactions include alkylation,2–4 silylation,4–7 bromination,3–4 acylation,2,4,8 a-aminoalkylation 3,9 and a-hydroxyalkylation.2,10 The reaction of an a-lithiated phosphinimine with an aryl nitrile yielded, after hydrolysis, the iminophosphoranylenamine, which when further hydrolysed by a solution of sulfuric acid led to the b-ketophosphine oxide (Scheme 1);2 but no attempt was made to separate the lithiated intermediate.We report here our investigations of the addition reactions of Li{CH(R9)P(R)2]] NSiMe3} (R = Me, R9 = SiMe3; R = Ph, R9 = SiMe3; or R = Ph, R9 = H) and benzonitrile, the conversion of one of these adducts into the potassium analogue and crystal structural characterisation of [K{N(H)C(Ph)C(H)P(Ph)2]] NSiMe3}(tmen)]2 (tmen = Me2- NCH2CH2NMe2). Results and Discussion The synthesis and reactions of complexes Li{N(R9)C(Ph)C(H)P(R)2]] NSiMe3} (R = Me, R9 = SiMe3 1; R = Ph, R9 = SiMe3 2; or R = Ph, R9 = H 3) are summarised in Scheme 2.Treatment of the appropriate compound Li{CH(R9)P(R)2]] NSiMe3} [obtained in situ from the phosphinimine CH2(R9)P(R)2]] NSiMe3 and LiBun] with PhCN in Et2O proceeded readily at room temperature. Complexes 1–3 were formed as colourless crystals (1), pale yellow solid (3) or colourless oil (2). A 1,3-trimethylsilyl (when R9 = SiMe3) or hydrogen (when R9 = H) C æÆ N migration was observed in these reactions.Mixing of 2 with KOBut in hexane afforded a white precipitate of the potassium analogue 4. Crystallisation of 4 with tmen in hexane produced colourless crystals of the tmen adduct 5. From the mother-liquor of 5, a few colourless crystals were * E-Mail: m.f.lappert@sussex.ac.uk further separated; a single-crystal X-ray diffraction study showed that these were those of another potassium complex 6.Formation of the very small amount of complex 6 is assumed to have been the result of partial hydrolysis of Scheme 1 Reagents and conditions: (i) LiNPri 2 or LiBun, 230 8C, tetrahydrofuran (thf); (ii) R9CN (R9 = Ph or 4-MeC6H4); (iii) H2O; (iv) H2O–H2SO4 R Ph2 P NPh R Ph2 P NPh R Ph2 P O R¢ NH2 R¢ O (i ), (ii ), (iii ) (iv ) R = H or Me Scheme 2 Reagents and conditions: (i) PhCN, Et2O, 250 8C to ca. 20 8C, 3 h; (ii) KOBut, hexane, ca. 20 8C, 3 h; (iii) tmen, hexane, ca. 20 8C; (iv) H2O, thf, 220 8C to ca. 20 8C, 1 h N R2P CH Li R2P Me3SiN H NSiMe3 SiMe3 R¢ Ph 7 R = Me 8 R = Ph R = Me, R¢ = SiMe3 R = Ph, R¢ = SiMe3 R = Ph, R¢ = H Ph2P Me3SiN K NSiMe3 Ph Ph2P Me3SiN K NH Ph Ph2P Me3SiN K NSiMe3 Ph R2P Me3SiN Li NR¢ Ph tmen tmen 5 6 4 1 R = Me, R¢ = SiMe3 2 R = Ph, R¢ = SiMe3 (not isolated) 3 R = Ph, R¢ = H (i ) (iv ) R¢ = SiMe3 (ii ) R = Ph R¢ = SiMe3 + (iii )1954 J. Chem. Soc., Dalton Trans., 1997, Pages 1953–1956 Table 1 Selected intramolecular distances (Å) and angles (8) for complex 6 K(1)]N(1) K(1)]N(3) K(1)]N(5) K(1)]C(2) K(1)]C(25) K(2)]N(2) N(2)]K(1)]N(5) N(5)]K(1)]N(4) N(5)]K(1)]N(3) N(2)]K(1)]C(2) N(4)]K(1)]C(2) N(2)]K(1)]C(25) N(4)]K(1)]C(25) C(2)]K(1)]C(25) N(5)]K(1)]N(1) N(3)]K(1)]N(1) C(25)]K(1)]N(1) N(5)]K(1)]C(1) N(3)]K(1)]C(1) C(25)]K(1)]C(1) N(2)]K(1)]C(24) N(4)]K(1)]C(24) 3.137(8) 3.033(8) 2.973(11) 3.072(9) 3.116(9) 2.858(8) 144.4(3) 137.8(3) 99.0(3) 25.0(2) 102.8(2) 102.4(2) 24.7(2) 127.4(3) 106.0(3) 154.4(2) 97.7(2) 102.5(3) 127.3(2) 139.3(3) 120.0(2) 45.7(2) K(1)]N(2) K(1)]N(4) K(1)]C(1) K(1)]C(24) K(2)]N(1) K(2)]N(4) N(2)]K(1)]N(4) N(2)]K(1)]N(3) N(4)]K(1)]N(3) N(5)]K(1)]C(2) N(3)]K(1)]C(2) N(5)]K(1)]C(25) N(3)]K(1)]C(25) N(2)]K(1)]N(1) N(4)]K(1)]N(1) C(2)]K(1)]N(1) N(2)]K(1)]C(1) N(4)]K(1)]C(1) C(2)]K(1)]C(1) N(1)]K(1)]C(1) N(5)]K(1)]C(24) N(3)]K(1)]C(24) 2.874(9) 3.020(8) 3.209(10) 3.251(10) 2.978(8) 2.745(8) 77.7(2) 93.1(2) 67.4(2) 119.4(3) 102.8(2) 113.2(3) 66.7(2) 69.7(2) 89.9(2) 69.7(2) 45.6(2) 117.7(2) 26.0(2) 52/4(2) 93.5(3) 52.8(2) K(2)]N(7) P(1)]N(1) P(2)]N(3) Si(2)]N(3) N(2)]C(2) C(1)]C(2) C(2)]K(1)]C(24) N(1)]K(1)]C(24) N(4)]K(2)]N(8) N(8)]K(2)]N(2) N(8)]K(2)]N(1) N(4)]K(2)]N(7) N(2)]K(2)]N(7) P(1)]N(1)]Si(1) Si(1)]N(1)]K(2) Si(1)]N(1)]K(1) C(2)]N(2)]K(2) K(2)]N(2)]K(1) P(2)]N(3)]K(1) C(25)]N(4)]K(2) K(2)]N(4)]K(1) N(2)]C(2)]C(1) 3.028(10) 1.568(8) 1.593(8) 1.669(8) 1.302(11) 1.418(12) 143.3(2) 118.9(2) 99.2(3) 87.3(3) 150.6(3) 114.0(3) 145.5(3) 134.6(5) 108.8(4) 121.5(4) 137.7(6) 82.6(2) 99.8(3) 132.7(6) 81.9(2) 122.0(8) K(2)]N(8) P(1)]C(1) P(2)]C(24) Si(1)]N(1) N(4)]C(25) C(24)]C(25) C(25)]K(1)]C(24) C(1)]K(1)]C(24) N(4)]K(2)]N(2) N(4)]K(2)]N(1) N(2)]K(2)]N(1) N(8)]K(2)]N(7) N(1)]K(2)]N(7) P(1)]N(1)]K(2) P(1)]N(1)]K(1) K(2)]N(1)]K(1) C(2)]N(2)]K(1) P(2)]N(3)]Si(2) Si(2)]N(3)]K(1) C(25)]N(4)]K(1) C(2)]C(1)]P(1) C(25)]C(24)]P(2) 2.852(11) 1.712(9) 1.741(9) 1.701(8) 1.315(11) 1.368(12) 24.7(2) 163.3(3) 82.7(2) 98.9(2) 72.3(2) 61.2(3) 129.6(3) 101.1(3) 97.9(4) 76.4(2) 86.0(6) 134.1(5) 126.0(4) 81.8(5) 122.9(7) 126.1(7) complex 5.The anionic ligand of 6 is identical to that in 3. However, attempts to prepare 6 by reaction of 3 with KOBut in hexane or benzene and then to crystallise it as a tmen complex were unsuccessful; instead, an unidentified yellow, crystalline compound was obtained. Treatment of complex 1 or 2 with 1 equivalent of H2O in thf gave in high yield N(SiMe3)C(Ph)C(H)P(R)2N(SiMe3)H as colourless crystals 7 (R = Me) or pale yellow oil 8 (R = Ph).The oily complex 2 was characterised solely by its conversion into the potassium analogue 4, while each of the crystalline complexes 1 and 3 gave satisfactory microanalytical data as well as 1H, 13C-{1H}, 31P-{1H} and 7Li-{1H} NMR spectra. The 1H NMR spectrum of Li{N(SiMe3)C(Ph)C(H)P(Me)2]] NSiMe3} 1 showed two sets of broad signals for the phosphorus-bound methyl groups.A variable-temperature 1H NMR study showed that the two sets of signals observed at 291 K coalesced to a sharp singlet when the temperature was raised to 348 K. This is ascribed to the fact that at the lower temperature the lithium atom is firmly attached to the ligand giving rise to distinct axial and equatorial P-methyl groups, while at the higher temperature rapid inversion occurs. The 13C-{1H} NMR spectrum of 1 provided evidence for the migration of the trimethylsilyl group.In the 13C-{1H} NMR spectra of Li{CH(SiMe3)P(Ph)2]] NSiMe3} or K{CH(SiMe3)P(Me)2]] NSi- Me3}, P]SiMe3 coupling was observed for both trimethylsilyl groups, each appearing as a doublet.11 However, in the 13C-{1H} NMR spectrum of 1, only one silylmethyl signal was a doublet, the other a singlet. Such 1,3-trimethylsilyl migrations from carbon to nitrogen have previously been observed in the reaction of MCH(R)SiMe3 (M = Li or K; R = Ph or SiMe3) with a nitrile (ButCN or PhCN), giving rise to M{N(SiMe3)C(But)CHR} 12 or M{N(SiMe3)C(Ph)C(H)C(Ph)NSiMe3}.13 However, for Li{CH(SiMe3)P(R)2]] NSiMe3}, no reaction with ButCN occurred, even under reflux in thf; this is attributed to a steric effect.The 1H NMR spectrum of Li{N(H)C(Ph)C(H)P(Ph)2]] NP N N Li Ph Me Me Me3Si SiMe3 1 SiMe3} 3 showed two equal intensity signals at d 4.16 and d 4.43. The former was a doublet, J(1H]31P) 24.3 Hz, and the latter a singlet (NH). This clearly shows that one of the CH2 protons had migrated from carbon to nitrogen in the course of the reaction of Li{CH2P(Ph)2]] NSiMe3} with PhCN.The 13C-{1H} NMR spectrum showed the silylmethyl signal as a doublet. This imples that the trimethylsilyl group lies on the nitrogen atom attached to the phosphorus, rather than on the more remote nitrogen atom. The 1H, 13C-{1H} and 31P-{1H} NMR spectra of the potassium compounds 4 and 5 were consistent with the structures shown in Scheme 2. Their 13C-{H} NMR spectra showed a similar feature to that in complex 1, in that only one silylmethyl was split by 31P coupling.Compounds 7 and 8 are neutral iminophosphoranylenamines. Their 1H NMR spectra showed that the NH signals appeared at d 8.39 (7) and 8.29 (8). The high frequency chemical shifts indicate the presence of hydrogen bonds in both molecules. The molecular structure and the atom numbering scheme of [K{N(H)C(Ph)C(H)P(Ph)2]] NSiMe3}(tmen)]2 6 are shown in Fig. 1. Selected bond distances and angles are presented in Table 1.Fig. 1 Molecular structure of [K{N(H)C(Ph)C(H)P(Ph)2]] NSiMe3}- (tmen)]2 6 with the atom numbering schemeJ. Chem. Soc., Dalton Trans., 1997, Pages 1953–1956 1955 Crystalline 6 is a dinuclear molecule with each of the two potassium atoms in a different co-ordination environment. The atom K(1) has co-ordination number nine taking the N(H)C(Ph)C(H)-moiety to occupy three co-ordination sites and K(2) co-ordination number five. As a result of the very crowded environment around K(1), only one nitrogen atom of tmen is co-ordinated to this potasium atom.Unlike K(1), K(2) is bonded to a chelating tmen. It is also noteworthy that the ligand co-ordination is not h5, there being no P]K bond. It seems that there is a combination of h3-azaallyl and phosphinimine co-ordination, the latter bonding to the potassium through the free-electron pair on nitrogen. The K]C bond lengths are in the range 3.072(9)–3.251(10) Å, but the K]N bond lengths vary rather more widely, from 2.745(8) to 3.137(8) Å.The P atom has a slightly disorted tetrahedral geometry. The mean P]N bond length of 1.58 Å is indicative of a double bond.14–16 The P(1)]C(1) bond length [1.712(9) Å] is close to that of P(2)]C(24) [1.741(9) Å] and shows that the P]C bond is in between a formal single (1.83 Å)14,15 and double (1.57 Å)14 bond. There are no short K? ? ? H contacts in the vacant site trans to N(4). Intermolecular approach at this site is blocked by the phenyl group C(15)]C(20), and the shortest K(2) ? ? ? C contacts to this group are K(2) ? ? ? C(15) 3.48 and K(2) ? ? ? C(20) 3.45 Å.Experimental All reactions were performed under argon using standard Schlenk techniques. The thf and diethyl ether were dried using sodium–benzophenone; hexane and pentane were dried using sodium–potassium alloy. The NMR spectra were recorded on AC-P250, WM-360 or AMX-500 instruments, and the solvent resonances were used as the internal references for 1H and 13C spectra; H3PO4 (85% aqueous solution) and LiCl (1 mol dm23 aqueous solution) were the external references for 31P and 7Li NMR spectra, respectively.Infrared spectra were recorded on a Perkin-Elmer 1720 FT spectrometer as liquid films or Nujol mulls using KBr windows, and elemental analyses were carried out by Medac Ltd, Brunel University. Melting points were determined under argon in sealed capillaries on an electrothermal apparatus and were uncorrected.Preparations The phosphinimines CH2(R)P(Ph)2]] NSiMe3. (a) R = H. A mixture of methyl(diphenyl)phosphine (10.0 g, 0.05 mmol) and trimethylsilyl azide (6.6 cm3, 0.05 mmol) was heated at 130 8C for 6 h with stirring. The unreacted starting materials were removed at 25 8C (1025 Torr) leaving CH3P(Ph)2]] NSiMe3 (13.6 g, 95%), as a colourless oil. 1H NMR (298 K, CDCl3): d 0.02 (s, 9 H, SiMe3), 1.96 [d, 3 H, 2J(1H]31P) = 12.7 Hz, CH3], 7.46 (s, 6 H, Ph), 7.65–7.75 (m, 4 H, Ph).(b) R = SiMe3. A solution of LiCH2SiMe3 (50 cm3 of a 1 mol dm23 solution in pentane, 0.05 mol) was added dropwise to a solution diphenylphosphorus(III) chloride (11.0 g, 0.05 mol) in tetrahydrofuran (50 cm3) at 278 8C with stirring. The mixture was allowed to warm to room temperature with stirring during 15 h. The precipitate was filtered off and solvent was removed from the filtrate in vacuo. The residue was distilled to yield CH2(SiMe3)PPh2 (10.7 g, 78%), b.p. 99–101 8C (0.02 Torr). 1H NMR (298 K, CDCl3): d 0.01 (s, 9 H, SiMe3), 1.43 [d, 2 H, 2J(1H–31P) = 10.0 Hz, CH2], 7.34–7.36 (m, 6 H, Ph), 7.51–7.54 (m, 4 H, Ph). A mixture of CH2(SiMe3)PPh2 (10.7 g, 0.039 mol) and trimethylsilyl azide (5.2 cm3, 0.039 mol) was heated at 130 8C for 6 h with stirring. The unreacted starting materials were removed at 25 8C (1025 Torr), leaving a residue of CH2(SiMe3)P(Ph)2]] NSiMe3 (13.7 g, 97.5%). 1H NMR (298 K, CDCl3): d 20.02 (s, 9 H, SiMe3), 20.01 (s, 9 H, SiMe3), 1.69 [d, 2 H, 2J(1H]31P) = 15.3 Hz, CH2], 7.39–7.42 (m, 6 H, Ph), 7.64– 7.70 (m, 4 H, Ph).Li{N(SiMe3)C(Ph)C(H)P(Me)2]] NSiMe3} 1. Benzonitrile (0.58 cm3, 5.69 mmol) was added dropwise to a stirred solution of Li{CH(SiMe3)P(Me)2]] NSiMe3} [from CH2(SiMe3)P- (Me)2]] NSiMe3 16 (1.33 g, 5.68 mmol) and LiBun (3.6 cm3 of a 1.6 mol dm23 solution in hexane, 5.68 mmol)] (1.37 g, 5.68 mmol) in diethyl ether (30 cm3) at ca. 20 8C. The solution was stirred for 4 h and the solvent was then removed in vacuo.The resultant solid was recrystallised from pentane to yield the colourless crystalline compound 1 (1.4 g, 72%) (Found: C, 56.5; H, 8.95; N, 7.7. C16H30LiN2PSi2 requires C, 55.8; H, 8.8; N, 8.15%), m.p. 115–126 8C. NMR (298 K, C6D6): 1H, d 0.17 (s, 9 H, SiMe3); 0.34 (s, 9 H, SiMe3); 0.84–1.05 (br, 3 H, PMe); 1.42– 1.60 (br, 3 H, PMe); 3.69 (d, 1 H, CH, J = 26.3 Hz); 7.15 (s, 3 H), 7.47 (s, 2 H) (Ph). 13C-{1H}, d 3.80 (s), 4.06 (d, J = 3.5) (SiMe3); 16.21 (d, J = 45), 19.58 (d, J = 60) (PMe); 77.55 (d, CH, J = 118.1); 127.68 (s), 128.65 (s), 128.76 (s), 147.42 (d, J = 14.4 Hz) (Ph); 178.91 (s, PhC). 31P-{1H}, d 17.61. 7Li-{1H}, d 0.67. Li{N(H)C(Ph)C(H)P(Ph)2]] NSiMe3} 3. Benzonitrile (0.4 cm3, 3.92 mmol) was added dropwise at ca. 250 8C to a diethyl ether (ca. 30 cm3) solution of Li{CH2P(Ph)2]] NSiMe3}, prepared from CH3P(Ph)2]] NSiMe3 (1.09 g, 3.80 mmol) and LiBun (2.4 cm3 of a 1.6 mol dm23 solution in hexane, 3.84 mmol).The mixture was allowed to warm to room temperature and was stirred overnight, then filtered to remove some yellow precipitate. Volatiles were removed from the filtrate in vacuo and the solid residue was washed with pentane (2 × 15 cm3). Drying in vacuo yielded the free-flowing, pale yellow powder 3 (1.26 g, 84%) (Found: C, 69.7; H, 6.65; N, 6.95. C23H26LiN2PSi requires C, 69.7; H, 6.6; N, 7.05%), m.p. 206–208 8C. NMR (298 K, C6D6 1 10% [2H8]thf): 1H, d 0.00 (s, 9 H, SiMe3); 4.16 (d, 1 H, CH, J = 24.3 Hz); 4.43 (s, 1 H, NH); 7.08 (s), 7.22 (s), 7.79 (s), 7.84 (s) (Ph). 13C-{1H}, d 4.32 (d, SiMe3, J = 3.8); 67.93 (d, CH, J = 78.1); 126.47 (s), 128.07 (d, J = 11.3), 128.26 (s), 130.14 (d, J = 2.6), 132.44 (d, J = 9.9), 139.59 (d, J = 93.5), 148.53 (d, J = 11.8 Hz) (Ph); 175.88 (s, PhC). 31P-{1H}, d 14.84. 7Li-{1H}, d 0.68. K{N(SiMe3)C(Ph)C(H)P(Ph)2]] NSiMe3} 4. Benzonitrile (0.28 cm3, 2.75 mmol) was added dropwise to a stirred solution of Li{CH(SiMe3)P(Ph)2]] NSiMe3] (see synthesis of 1) (1.0 g, 2.74 mmol) in diethyl ether (30 cm3) at 250 8C.The solution was allowed to warm to room temperature and was stirred for 4 h. The solvent was removed in vacuo and the resultant redbrown oil was redissolved in hexane (30 cm3). Solid KOBut (0.31 g, 2.77 mmol) was added with stirring at room temperature. After further stirring for 3 h, the mixture was filtered. The white precipitate was washed with hexane (2 × 20 cm3), then dried in vacuo leaving the white complex 4 (1.15 g, 84%) (Found: C, 61.4; H, 6.8; N, 5.25.C26H34KN2PSi2 requires C, 62.4; H, 6.85; N, 5.6%), m.p. 174–177 8C. NMR (298 K, C5D5N): 1N, d 20.17 (s, 9 H, SiMe3); 0.12 (s, 9 H, SiMe3); 4.42 (d, 1 H, CH, J = 25.8 Hz); 7.18–7.57 (m), 8.14–8.19 (m) (Ph). 13C-{1H}, d 3.69 (s), 4.68 (d, J = 3.7) (SiMe3); 76.31 (d, CH, J = 129); 126.51 (s), 127.59 (d, J = 5.4), 127.98 (d, J = 11.0), 129.29 (s), 132.06 (d, J = 9.4), 141.59 (d, J = 95.1), 151.04 (d, J = 16.9 Hz) (Ph); 176.10 (s, PhC). 31P-{1H}, d 5.90. K{N(SiMe3)C(Ph)C(H)P(Ph)2]] NSiMe3}(tmen) 5. Tetramethylethylenediamine (0.9 cm3, 5.97 mmol) was added dropwise to a stirred suspension of 4 (1.2 g, 2.4 mmol) in hexane (ca. 30 cm3) at room temperature. The solid dissolved gradually, yielding a pale yellow solution, which was concentrated to ca. 3 cm3 and set aside for several hours at room temperature to yield colourless crystals of compound 5 (1.19 g, 80%) (Found: C, 61.9; H, 8.0; N, 9.1.C32H50KN4PSi2 requires C, 62.3; H, 8.15; N, 9.1%), m.p. 111–113 8C. NMR (298 K, C6D6): 1H, d 20.12 (s, 91956 J. Chem. Soc., Dalton Trans., 1997, Pages 1953–1956 H, SiMe3); 0.19 (s, 9 H, SiMe3); 1.91 (s, 12 H, tmen), 1.92 (s, 4 H, tmen); 4.40 (d, 1 H, CH, J = 25.7 Hz); 7.13–7.23 (m), 7.56– 7.59 (m), 7.99–8.05 (m) (Ph). 13C-{1H}, d 3.70 (s), 4.68 (d, J = 3.9 Hz) (SiMe3); 45.47 (s), 57.36 (s) (tmen); 76.38 (d, CH, J = 126.6); 126.94 (s), 127.68 (d, J = 6.1), 128.05 (d, J = 11.3), 129.49 (d, J = 2.6), 131.91 (d, J = 9.7), 140.84 (d, J = 95.7), 150.44 (d, J = 16.8) (Ph); 176.72 (d, PhC, J = 3.9 Hz). 31P-{1H}, d 7.02. N(SiMe3)C(Ph)C(H)P(Me)2N(SiMe3)H 7. Water (0.2 cm3, 11.11 mmol) was added at 220 8C with stirring to a solution of Li{N(SiMe3)C(Ph)C(H)P(Me)2]] NSiMe3} 1 (3.54 g, 10.03 mmol) in thf (20 cm3). The mixture was stirred for 30 min at room temperature. The thf was removed in vacuo and the residue was extracted with diethyl ether.The extract was filtered and the filtrate was concentrated in vacuo. Crystallisation of the residue from pentane yielded colourless crystals of 7 (2.87 g, 85%) (Found: C, 56.6; H, 9.15; N, 8.25. C16H31N2PSi2 requires C, 56.8; H, 9.25; N, 8.25%), m.p. 116–118 8C. NMR (298 K, CDCl3): 1H, d 20.08 (s, 9 H, SiMe3); 0.09 (s, 9 H, SiMe3); 1.45 (d, 6 H, PMe2, J = 12.5); 3.83 (d, 1 H, CH, J = 24.6 Hz); 7.29– 7.38 (m, Ph); 8.39 (s, 1 H, NH). 13C-{1H}, d 20.02 (s), 2.80 (d, J = 4.0) (SiMe3); 20.47 (d, PMe, J = 64.3); 86.74 (d, CH, J = 114.3); 126.53 (s), 126.59 (s), 127.20 (s), 140.45 (d, J = 15.1) (Ph); 162.81 (d, PhC, J = 10.0 Hz). 31P-{1H}, d 6.46.IR (Nujol): nmax(cm21) 3300–3650 (br) (NH), 1590s, 1609s (C]] C, Ph), 1251s (P]] N). N(SiMe3)C(Ph)C(H)P(Ph)2N(SiMe3)H 8. Similarly, as for 7, Li{N(SiMe3)C(Ph)C(H)P(Ph)2]] NSiMe3} 2, prepared from Li{CH(SiMe3)P(Ph)2]] NSiMe3} (1.32 g, 3.62 mmol) and PhCN (0.37 cm3, 3.63 mmol), was treated with H2O (0.07 cm3, 3.89 mmol) to yield the pale yellow oil 8 (1.46 g, 87%) (Found: C, 67.2; H, 7.5; N, 5.4.C26H35N2PSi2 requires C, 67.5; H, 7.6; N, 6.05%). NMR (298 K, CDCl3): 1H, d 20.08 (s, 9 H, SiMe3); 20.05 (s, 9 H, SiMe3); 4.32 (d, 1 H, CH, J = 23.2 Hz); 7.33–7.48 (m), 7.67–7.73 (m) (Ph); 8.29 (s, 1 H, NH). 13C-{1H}, d 0.00 (s), 2.51 (d, J = 3.9) (SiMe3); 85.77 (d, CH, J = 118.5); 126.96 (d, J = 4.4), 127.44 (s), 129.00 (s), 129.77 (d, J = 4.6), 131.17 (d, J = 19.5), 136.33 (d, J = 94.8), 140.45 (d, J = 15.0 Hz) (Ph); 163.83 (s, PhC). 31P-{1H}, d 3.99.IR (liquid film): nmax(cm21) 3300–3500 (br) (NH), 1567vs, 1588vs (C]] C, Ph), 1256vs (P]] N). Crystallography Crystallographic details are given in Table 2. Single crystals of complex 6 were mounted in Lindemann capillaries under argon. Data were collected on an Enraf-Nonius CAD4 diffractometer in the q–2q mode with monochromated Mo-Ka radiation (l = 0.710 73 Å). The structure was solved by direct methods (SHELXS 86) 17 and refined by full-matrix least squares on all F2 (SHELXL 93).18 All non-H atoms were anisotropic.The H atoms on N(2) and N(4) were located on a difference map and refined with Uiso(H) = 1.5 Ueq(N) and the N]H bond lengths restrained to 0.89 Å. Other H atoms were included in riding mode with Uiso(H) equal to 1.2 Ueq(C) or 1.5 Ueq(C) for methyl groups. Atomic coordinates, thermal parameters, and bond lengths and angles have been deposited at the Cambridge Crystallographic Data Centre (CCDC). See Instructions for Authors, J.Chem. Soc., Dalton Trans., 1997, Issue 1. Any request to the CCDC for this material should quote the full literature citation and the reference number 186/482. Acknowledgements We thank the Chinese Government and the British Council for the award of a studentship (to Z.-X. W.), the Leverhulme Trust for the award of an Emeritus Fellowship (to M. F. L.), Dr. F. Reed (FMC) for n-butyllithium and EPSRC for other support.References 1 H.-J. Cristau, Chem. Rev., 1994, 94, 1299. 2 J. Barluenga, F. Lopez and F. Palacios, J. Chem. Res., 1985, (S) 211; (M) 2541. 3 A. K. Roy, U. G. Wettermark, G. M. Scheide, P. Wisian-Neilson and R. H. Neilson, Phosphorus Sulfur Relat. Elem., 1987, 33, 147. 4 K. D. Gallicano, R. T. Oakley, N. L. Paddock and R. D. Sharma, Can. J. Chem., 1981, 59, 2654. 5 H. Schmidbaur and G. Jonas, Chem. Ber., 1967, 100, 1120. 6 U. G. Wettermark, P. Wisian-Neilson, G. M. Scheide and R.H. Neilson, Organometallics, 1987, 6, 959. 7 A. K. Roy, R. Hani, R. H. Neilson and P. Wisian-Neilson, Organometallics, 1987, 6, 378. 8 J. Barluenga, F. Lopez and F. Palacios, J. Chem. Soc., Chem. Commun., 1985, 1681. 9 S. O. Grim and P. B. Kettler, J. Chem. Soc., Chem. Commun., 1991, 979. 10 J. Barluenga, F. Lopez and F. Palacios, Synthesis, 1988, 562. 11 P. B. Hitchcock, M. F. Lappert and Z.-X. Wang, Chem. Commun., in the press. 12 P. B. Hitchcock, M. F. Lappert and D.-S. Liu, J. Chem. Soc., Chem. Commun., 1994, 2637. 13 P. B. Hitchcock, M. F. Lappert and D.-S. Liu, J. Chem. Soc., Chem. Commun., 1994, 1699. 14 See, for example, D. E. C. Corbridge, Phosphorus, Elsevier, Amsterdam, 1985. 15 P. Imhoff, R. van Asselt, C. J. Elsevier, K. Vrieze, K. Goubitz, K. F. van Malssen and C. H. Stam, Phosphorus, Sulfur Silicon Relat. Elem., 1990, 47, 401. 16 J. C. Wilburn and R. H. Neilson, Inorg. Chem., 1979, 18, 347. 17 G. M. Sheldrick, SHELXS 86, a package for the solution of crystal structures, University of Göttingen, 1985. 18 G. M. Sheldrick, SHELXL 93, a package for crystal structure refinement, University of Göttingen, 1993. Received 12th December 1996; Paper 6/08355B Table 2 Crystal data and refinement for complex 6 Formula M Crystal system Space group a/Å b/Å c/Å a/8 b/8 g/8 U/Å3 Z D/g cm23 Crystal size/mm Radiation l/Å F(000) m(Mo-Ka)/mm21 Temperature/K Total reflections Independent reflections Reflections with I > 2s(I) R1 [I > 2s(I)] a wR2 (all data) b C58H84K2N8P2Si2 1089.6 Triclinic P1 (no. 1) 11.310(3) 11.856(5) 13.263(4) 70.51(3) 70.83(3) 83.59(3) 1583.6(9) 1 1.14 0.2 × 0.2 × 0.2 0.710 73 584 0.28 293(2) 5553 5553 3319 0.067 0.174 a R1 = S||Fo| 2 |Fc||/S|Fo|. b wR2 = [Sw(Fo 2 2 Fc 2)2/SwFo 4]�

ISSN:1477-9226    

DOI:10.1039/a608355b

出版商:RSC    

年代:1997

数据来源: RSC

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DALTON J. Chem. Soc., Dalton Trans., 1997, Pages 1957–1960 1957 Matrix-isolation studies on the vaporisation of alkali-metal tellurites: the characterisation of molecular M2TeO3 species (M = K, Rb or Cs) † Trevor N. Day, Robin A. Gomme and J. Steven Ogden* Department of Chemistry, The University, Southampton SO17 1BJ, UK Solid samples of alkali-metal tellurites (M2TeO3; M = K, Rb or Cs) have been heated in vacuo and the vaporisation products studied by mass spectrometry and matrix-isolation infrared spectroscopy.The molecular species M2TeO3 are identified by mass spectrometry as important products of vaporisation, and subsequent matrix-isolation studies, which incorporated oxygen-18 data, established both vibrational assignments and probable structures for these species. In particular, molecular K2TeO3 in argon has prominent IR absorptions at ca. 752, 742, 672, 360 and 306 cm21 and a structure based on a pyramidal TeO3 unit. Alkali-metal tellurites (M2TeO3, where M = Cs, Rb or K) have been studied extensively in the solid phase.The crystal structures of rubidium tellurite,1 potassium tellurite trihydrate,2 caesium tellurite3 and potassium tellurite 4 have been determined. Their structures all contain TeO3 22 units, whereas other tellurium(IV) compounds, e.g. a-TeO2, involve four oxygen atoms co-ordinating to each tellurium atom.5 Studies concerned with heating alkali-metal tellurites in air have shown partial oxidation of TeIV to TeVI between 500 and 550 8C, a process which is reversed above 800 8C.6,7 The salts were found to be stable up to 1000 8C in air.Two Knudsen-cell mass spectrometric studies of the vaporisation of caesium tellurite have been reported. The first of these 8 established the existence of Cs2TeO3 as a probable molecular species, and obtained the standard enthalpy of vaporisation, whilst a more recent paper 9 established the congruent vaporisation of caesium tellurite and also studied the thermal dissociation of caesium tellurate.A calculation of the entropy of the gaseous ion TeO3 22 has been reported 10 based on data taken from the solid phase. The analogous alkali-metal selenites have previously been studied by matrix isolation and mass spectrometry,11 where the predominant vapour phase species were found to be M2SeO3, with structures based on the pyramidal SeO3 22 unit. In this study the vapour-phase species above the alkali-metal tellurites have been similarly investigated and the basic vibrational data for one particular system (K2TeO3) supplemented by experiments on the 18O-enriched material in order to assist in the characterisation.Experimental The alkali-metal tellurites used in these studies were obtained by two separate routes. The majority of samples were prepared in situ by heating an equimolar mixture of the alkali-metal carbonate with tellurium dioxide in vacuo at ca. 800 8C according to the recommended procedure.12 These samples were subsequently reground and heated in a platinum boat located inside an alumina tube which was heated inductively using a tantalum susceptor.Vaporisation took place in vacuo at temperatures typically in excess of ª 900 8C. Samples of potassium tellurite were also obtained commercially (Aldrich, as hydrate) and vaporised in the same way, giving identical results. Oxygen-18-enriched potassium tellurite was prepared by dissolving the commercial sample in an excess of H2 18O (90 atom %) and leaving for 10 d.At the end of this † Non-SI unit employed: dyn = 1025 N. period the water was removed by vacuum distillation and the sample vaporised as described above. The matrix-isolation apparatus and general methodology used for these experiments have been described elsewhere.11,13 Both nitrogen and argon (BOC, 99.999%) were used as matrix gases in an estimated >1000-fold excess, and infrared spectra were recorded using a Perkin-Elmer PE983G instrument.Controlled-diffusion studies were carried out up to ca. 35 K. Mass spectra were obtained using the VG quadrupole models SXP600 and 12-12S operating in cross-beam mode. Results Our mass spectrometric studies on the vaporisation of alkalimetal tellurites were in good agreement with previous work.8,9 For the caesium salt we observed a very intense signal corresponding to Cs1, but a prominent signal was also obtained for the Cs2TeO3 1, together with weaker features corresponding to Cs2TeO2 1, TeO1 and Te1, which were assigned as fragments by varying the electron energy.The rubidium and potassium systems similarly yielded parent ions M2TeO3 1 and related fragments. Matrix-isolation studies generally involved deposition times of up to 2 h, and during the initial stages IR spectra typically showed the characteristic absorptions of CO2, notably at ca. 662 cm21, which corresponds to the bending mode. However, the evolution of CO2 decreased with time, with the result that its absorptions could be effectively removed from subsequent spectra by subtraction.Fig. 1 shows part of the nitrogen-matrix IR spectrum (800– 250 cm21) obtained from the vaporisation of potassium tellurite, after removal of the CO2 absorption. Two prominent bands are observed at 740.8 and 672.8 cm21 and these are accompanied by weaker features at 750.9, 724.2, 713 (sh), 360 and 306 Fig. 1 Infrared spectrum of molecular K2TeO3 isolated in a nitrogen matrix.The feature denoted P is assigned to a polymer1958 J. Chem. Soc., Dalton Trans., 1997, Pages 1957–1960 cm21. Although all these bands showed essentially the same growth behaviour on deposition, controlled diffusion to ca. 35 K resulted in the continued growth of the bands at 724.2 and 713 cm21, but a decrease in the remaining five features. These latter bands at 750.9, 740.8, 672.8, 360 and 306 cm21 are therefore assigned to a species A, which is chemically distinct from a second species B with absorptions at 724.2 and 713 cm21.On the basis of this diffusion behaviour species B is provisionally assigned as ‘polymeric’, whilst species A is identified as a possible monomer. It may be significant, however, that in this study or in the previous mass spectrometric reports 8,9 signifi- cant amounts of polymer were not detected in the vapour phase. Corresponding features showing similar diffusion behaviour were also obtained from the rubidium and caesium systems and Table 1 summarises their band positions.Fig. 2(a) shows the higher-frequency spectral region obtained from K2TeO3 isolated in argon. This spectrum is very similar to that in nitrogen and the frequencies are included in Table 1. From these results it is evident that varying either the cation or the matrix environment produces only small shifts in vibrational frequency and we therefore assign all these features as vibrations involving essentially only Te/O motion in which there is little interaction with the matrix cage.Fig. 3(a) shows the argon-matrix spectrum in the higherfrequency region obtained from the vapour above 18O-enriched potassium tellurite. The CO2 bands have not been subtracted out, but comparison with Fig. 2(a) reveals a number of additional peaks, notably at ca. 710 and 640 cm21 and these are assigned to 18O-enriched ‘monomer’. Absorptions in the lowerfrequency region of this spectrum were too weak to reveal well defined maxima.Spectral Interpretation and Discussion On the basis of the mass spectrometric studies, in which the M2TeO3 1 peaks are assigned as molecular ions, we provisionally identify the ‘monomer’ species A as matrix-isolated molecular tellurites M2TeO3. The frequencies of the three high-frequency bands (Table 1) lie in the Te]O stretching region, and may be compared to those in the pyramidal TeO3 22 ion. In aqueous Fig. 2 Observed and calculated spectra in the Te]O stretching region for molecular K2TeO3 isolated in an argon matrix solution,14 this ion has C3v symmetry.The stretching modes transform as A1 1 E, with the A1 stretch lying at 748 cm21, close to the highest-frequency mode of species A, and the E at 703 cm21, almost exactly between the more intense features at ca. 740 and 670 cm21. The general appearance of this spectral region is very similar to that previously observed for alkali-metal selenites,11 and we propose a similar spectral interpretation for the tellurite spectra obtained here.Basically, we consider that coordination of the pyramidal TeO3 unit by two cations reduces the symmetry to Cs, with a resulting splitting of the degenerate E mode into two components, A9 1 A0. The A1 mode in the C3v ion becomes A9 in Cs, and as it is unlikely that two modes of the same symmetry end up close together in frequency, we assign the observed bands at ca. 750, 740 and 670 cm21 as A9, A0 and A9 respectively.This symmetry ordering, in which the parent E mode splits such that A0 > A9, predicts that the lowering of symmetry is such as to produce one Te]O bond which has a lower principal force constant than the other two. Fig. 4 shows three possible structures for molecular M2TeO3 species, all of which have Cs symmetry. The bis(monodentate) structure (a) has one terminal Te-O bond, which would be expected to be shorter and stronger than the two other Fig. 3 Observed and calculated argon-matrix spectra for ca. 25% 18Oenriched K2TeO3 species. The band denoted (*) is assigned to CO2. Calculated spectra are displayed at bandwidths of 0.2 and 2.5 cm21 Table 1 Vibration wavenumbers (cm21) observed in matrix-isolation studies on the vaporisation of alkali-metal tellurites K2TeO3 Rb2TeO3 Cs2TeO3 Species N2 Ar N2 N2 A B 750.9 740.8 672.8 360 307 724.2 713 (sh) 752.0 741.9 672.0 360 306 724.9 714 750.3 740.0 672.2 354 347 (sh) 303 725.5 713.4 748.9 736.4 673.6 350 340 (sh) 316 721.6 714.7 Accuracy: above 600 cm21, ± 0.5 cm21; below 600 cm21, ± 2 cm21.J. Chem.Soc., Dalton Trans., 1997, Pages 1957–1960 1959 Te]O bonds which link to the cations, and is therefore unlikely to fit the observed spectrum. However, in the other two structures (b) and (c) the unique Te]O bond could be considered to be weaker than the other two. In structure (b) the unique bond bridges to a single cation and would be expected to have more Fig. 4 Possible structures for molecular alkali-metal tellurites: (a) bis(monodentate)model, (b) mono-/bi-dentate model, (c) bis(bidentate) model Table 2 Observed and calculated vibration wavenumbers (cm21) for 18O-enriched K2TeO3 species isolated in argon matrices Observed Calculated (intensity) Assignment 752.0 747.6 741.9 709.9 706.6 (sh) 672.0 640.1 752.0 (1.1) 748.9 (0.46) 747.6 (1.34) 745.6 (0.51) 741.9 (3.6) 741.9 (1.2) 717.2 (0.1) 710.9 (1.57) 710.9 (0.04) 708.1 (0.53) 705.6 (0.36) 705.6 (0.12) 672.0 (3.57) 669.6 (2.48) 667.4 (0.42) 640.4 (1.04) 639.7 (0.71) 639.1 (0.12) A9 16OTe16O2 A9 18OTe16O2 A 16OTe16O18O A 18OTe16O18O A0 16OTe16O2 A0 18OTe16O2 A9 16OTe18O2 A 16OTe16O18O A9 18OTe18O2 A 18OTe16O18O A016OTe18O2 A018OTe18O2 A9 16OTe16O2 A 16OTe16O18O A9 16OTe18O2 A9 18OTe16O2 A 18OTe16O18O A9 18OTe18O2 Frequency calculations assume a Cs pyramidal structure, with one unique angle (108 8) and two angles of 115 8 as shown above.The value of the unique stretching constant Fd is taken as 3.90 mdyn Å21, and the other force constants have values Fr 4.65, Frd 0.41 and Frr 0.20 mdyn Å21.Relative band intensities (in parentheses) assume randomised 25% 18O enrichment. single-bond character than the other two, whilst in the bis- (bidentate) model (c) the unique O atom is close to two cations. Fig. 2(b) shows a spectral simulation of the Te]O stretching region for K2TeO3 isolated in argon using a ‘stretch only’ force field.11 The agreement is very satisfactory, but even this simple model requires a minimum of four independent stretching constants and the (arbitrary) choice of two angle parameters.Not surprisingly, there are a number of sets of such parameters which will generate a satisfactory spectral fit. However, incorporation of the oxygen-18 data introduces several additional frequency constraints and Table 2 summarises our final results for the simulation of the vibrations of K2TeO3 isolated in argon.In this model the basic structural unit appears as a distorted pyramid with one significantly small angle (108 8) and where the principal stretching constant of the unique bond (Fd) is smaller than for the other two. Fig. 3(b) shows the spectral fit for the 18O-enriched system using the parameters in Table 2, and the same parameters are employed for the simulation in Fig. 2(b). In addition to neglecting any coupling with other vibrations, this ‘stretch-only’ model for the Te]O modes neglects the effect of anharmonicity, and in achieving a satisfactory spectral fit these uncertainties can be expected to be reflected as errors not only in the values of the force constants, but in the O]Te]O bond angles.Nevertheless, we believe that the prediction of one relatively small angle and two wider ones is correct and we note that this is also the model proposed for the analogous M2SeO3 species.11 Calculations are currently in progress in an attempt to provide a theoretical justification for distinguishing between models (b) and (c) in Fig. 4 for both tellurites and selenites.15 The final assignment to be made concerns the spectral region 300–400 cm21. For K2TeO3 two bands are observed in this region, at 360 and 306 cm21, and for the rubidium and caesium salts similar features occur, with associated shoulders (Table 1). The TeO3 22 ion in aqueous solution has bending modes at 364 (A1) and 326 (E) cm21, and we provisionally assign the band at 360 cm21 as A9 (correlating with A1 in C3v) and that at 306 cm21 as one (or both) of the components of the split E mode.This assignment is supported by the relative intensities of these two bands. In both the selenite system 11 and the related iodate system,16 the A1 XO3 bending modes are found to be higher in frequency and more intense than the E. Conclusion These experiments first confirm that vaporisation of alkalimetal tellurites leads to the formation of molecular M2TeO3 and that the mass spectrometric studies are indeed looking at the parent ion.The matrix-isolation studies establish a number of fundamental vibrations, which are assigned on the basis of a distorted pyramidal structure for the TeO3 unit. A detailed vibrational assignment concludes that in this structure one Te]O bond has a lower stretching constant than the other two, and two possible models are proposed which satisfy this condition. A second species was also identified in these matrix studies, which is believed to be a polymer.This species appears to be readily formed during controlled-diffusion studies and its presence in the matrix almost certainly occurs via limited aggregation during cocondensation, rather than as the result of trapping from the vapour phase. Acknowledgements We gratefully acknowledge the financial support of the EEC for this work and of AEA Technology, Winfrith, UK through the provision of research studentships (to T.N. D. and R. A. G.). We also wish to thank Dr. B. R. Bowsher and Dr. S. Dickinson for helpful discussions.1960 J. Chem. Soc., Dalton Trans., 1997, Pages 1957–1960 References 1 H. Thummel and R. Hoppe, Z. Naturforsch, Teil B, 1974, 29, 28. 2 G. B. Johansson and O. Lindqvist, Acta Crystallogr., Sect. B, 1978, 34, 2959. 3 B. O. Loopstra and K. Goubitz, Acta Crystallogr., Sect. C, 1986, 42, 520. 4 L. Andersen, V. Langer, A. Stromberg and D. Stromberg, Acta Crystallogr., Sect. B, 1989, 45, 344. 5 A. Stromberg, U. Wahlgren and O. Lindqvist, Chem. Phys., 1985, 100, 229. 6 K. K. Samplavska, T. A. Khachaturyan and M. K. Karapetyants, Redk. Shchelochnye Elem., 1969, 73. 7 T. A. Khachaturyan, K. K. Samplavskaya and M. K. Karapetyants, Izv. Vyssh. Uchecbn. Zaved., Khim. Khim. Tekhnol., 1969, 12, 855. 8 R. Portman, M. J. Quinn, N. H. Sagert, P. P. S. Saluja and D. H. Wren, Thermochim. Acta, 1989, 144, 21. 9 G. A. Semenov, L. A. Fokina and R. A. Mouldagalieva, J. Nucl. Mater., 1994, 210, 167. 10 A. Loewenschuss and Y. Marcus, Chem. Rev., 1984, 84, 89. 11 A. K. Brisdon, R. A.Gomme and J. S. Ogden, J. Phys. Chem., 1991, 95, 2927. 12 E. H. P. Cordfunke and V. M. Smit-Groen, Thermochim. Acta, 1984, 80, 181. 13 J. S. Ogden and R. S. Wyatt, J. Chem. Soc., Dalton Trans., 1987, 859. 14 H. Z. Siebert, Z. Anorg. Allg. Chem., 1955, 275, 225. 15 A. V. Marenich, T. P. Pogrebnaya, V. V. Sliznev and V. G. Solomonik, International Symposium Computer Assistance to Chemical Research, Russian Academy of Sciences, Moscow, December 1996. 16 K. A. Biggs, R. A. Gomme, J. T. Graham and J. S. Ogden, J. Phys. Chem., 1992, 96, 9738. Received 21st January 1997; Paper 7/00478H

ISSN:1477-9226    

DOI:10.1039/a700478h

出版商:RSC    

年代:1997

数据来源: RSC

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DALTON J. Chem. Soc., Dalton Trans., 1997, Pages 1961–1965 1961 Intercalation of carboxymethyl-‚-cyclodextrin into magnesium– aluminum layered double hydroxide † Hongting Zhao and George F. Vance * Soil and Environmental Chemistry Group, Department of Plant, Soil and Insect Sciences, University of Wyoming, Laramie, WY 82071-3354, USA New intercalated compounds have been prepared by incorporating carboxymethyl-b-cyclodextrin (CMCD) into magnesium–aluminum layered double hydroxide (Mg/Al LDH).The sorptive uptake of CMCD(3) and CMCD(14), with a degree of carboxymethyl substitution of 3 and 14, respectively, by Mg/Al LDH was examined at 65 8C, and the resulting complexes were characterized using X-ray diffraction (XRD) and FT-IR spectroscopy. Results indicated that Mg/Al LDH retained approximately twice as much CMCD(3) as compared to CMCD(14). Results of XRD and FT-IR spectroscopy confirmed that both CMCD(14) and CMCD(3) could be intercalated into Mg/Al LDH interlayers due to an ion-exchange process. After exposure to CMCD(14) and CMCD(3), the d-spacing of Mg/Al LDH expanded from 8.74 to 15.48 and 20.63 Å, respectively.Intercalated CMCD(14) molecules formed monolayer coverage in the Mg/Al LDH interlayers, with the cavity axis perpendicular to the LDH layer, while CMCD(3) molecules adopted either a parallel-monolayered arrangement or a perpendicularbilayer- like coverage. Both CMCD(14) and CMCD(3) molecules were believed to be loosely packed within the Mg/Al LDH interlayers.Layered double hydroxides (LDHs), also known as anionic clays, are an important class of materials currently receiving considerable attention in view of their potential technological importance as catalysts, ion exchangers, optical hosts, ceramic precursors and antacids.1 They are considered antitypes of 2/1 clay minerals (three-layered clay minerals) and consist of positively charged metal oxide/hydroxide sheets with intercalated anions and water molecules.At present, they are the only layered materials with positively charged layers, in contrast to a large variety of materials with negative layers. They have considerably higher ion-exchange capacities than do smectitic clays. The general chemical composition of LDHs is:1 [M12xMIII x- (OH)2]x1Xn2 x/n?zH2O, where M can be Li1 or one of several divalent cations (Ca21, Mg21, Ni21, Co21, Zn21, Mn21 or Cu21), MIII is a trivalent cation (Al31, Cr31, Co31, Ni31, Mn31 or Fe31), and X is an interlayer anion (e.g., Cl2, NO3 2, ClO4 2, CO3 22, SO4 22, or other inorganic anion).The interlayer charge is x = x for M = divalent cations and x = 2x 2 1 for M = monovalent cations. The M2 :M31 ratio is usually between 1 and 5. With positively charged LDH structures, several anionic species have been intercalated into the gallery region of LDH layers, with the resulting intercalates used for several applications in various fields. These include anionic surfactants, metallophthalocyanine tetrasulfonates, and silicate anions,1–4 as well as large polyoxometalates 5–8 and various organic acids and polymers.9–12 Cyclodextrins are cyclic oligosaccharides of D-glucopyranose that possess a unique non-polar cylindrical cavity.They are typical ‘host’ molecules that can trap a great variety of molecules having the size of one or two benzene rings, or even larger compounds possessing a side chain of comparable size, to form crystalline inclusion complexes.13 Cyclodextrins and their derivatives are very useful as micro-encapsulating agents for stabilizing volatile or toxic organic compounds.These compounds have found prospective applications in the food industry, pharmaceutical industry, production of organic chemicals, agriculture, cosmetics, tobacco industry, environmental protection, and in various other fields.13–15 Modified b-cyclodextrins have been incorporated into montmorillonite,16–19 a- and g- † Non-SI unit employed: M = mol dm23. zirconium phosphates,20–22 as well as zeolitic structures.23 However, to our knowledge, no information is currently available on the intercalation of cyclodextrins into LDH structures.Because carboxymethyl-b-cyclodextrin (CMCD) (in the form of a sodium salt) is a derivative of b-cyclodextrin and could easily disassociate into its negatively-charged species, we hypothesize that CMCD molecules may be sorbed into the interlayers of LDH structures.The inclusion of cyclodextrin molecules as interlayer guest species in LDHs would appear to be intriguing because the host–guest interaction could impart unique structural features and physicochemical properties to the complex. Cyclodextrin–LDH complexes would also appear to be a novel use of LDHs as well as cyclodextrin molecules. The objective of this study was, therefore, to develop a novel family of inclusion compounds by examining the uptake and intercalation properties of carboxymethyl-b-cyclodextrin by magnesium aluminum hydroxide (Mg/Al LDH), and to explore potential uses of LDHs and cyclodextrins in catalysis reactions and in adsorption technology.Experimental Preparation of layered double hydroxide Magnesium–aluminum layered double hydroxide with an ideal formula 1 of [Mg3Al(OH)8]NO3?2H2O was selected for our intercalation studies. Nitrate was used as the charge balancing anion because carbonate anions are often difficult to exchange in Mg/Al, Ca/Al and Ni/Al hydroxides.1,24 The Mg/Al LDH was prepared by a procedure similar to that described by Meyn et al.1 A solution of Mg(NO3)2?6H2O (64 g, 0.250 mol) and Al(NO3)3?9H2O (23.4 g, 0.125 mol) in distilled, deionized water (250 cm3) was added dropwise over 1 h to a solution of NaOH (25.0 g, 0.62 mol) and NaNO3 (36.4 g, 0.42 mol) in water (290 cm3).The mixture was held at 65 8C for 16 d. The precipitate was separated by centrifugation, dialysed and dried at 40 8C.Reagents Reagent grade, hydrated metal nitrates were purchased from Aldrich (Milwaukee, WI) or Sigma (St. Louis, MO) and used as1962 J. Chem. Soc., Dalton Trans., 1997, Pages 1961–1965 received. Carboxymethyl-b-cyclodextrins (CMCDs) with different substitution degrees (3 and 14) were generously supplied by the American Cerestar Company (Hammond, IN) and used without further purification. The CMCD has an empirical formula of (C42H702nO35)?(CH2COONa)n. The average degree of carboxymethyl substitution (n) is 3 for CMCD(3) and 14 for CMCD(14).The average molecular weight for CMCD(3) and CMCD(14) is 1375 and 2255, respectively. Sorption/desorption isotherms Isotherm sorption experiments were conducted using the batch method. The LDHs (100 mg) were weighed into screw-top Corex glass centrifuge tubes that were filled with aqueous CMCD solutions (25 cm3) with concentrations ranging from 5.95 × 1022 to 3.57 mM for CMCD(14) and from 5.95 × 1022 to 12.2 mM for CMCD(3).The suspensions were shaken intermittently at 65 8C for 24 h, centrifuged, and the cyclodextrin concentrations in the supernatant solutions measured using a Shimadzu TOC-5000 (Total Organic Carbon) Analyzer. Two sets of the adsorption experiments were carried out simultaneously. After centrifugation, one set was used for desorption studies to examine the reversibility of cyclodextrin sorption, and the other one was washed with water extensively, centrifuged, and dried at 40 8C for X-ray diffraction (XRD) and FT-IR analysis.The desorption studies were conducted using the dilution method as follows: after the sorption experiments, some of the equilibrated cyclodextrin solution (10–20 cm3) was removed and distilled, deionized water (10–20 cm3) added. Samples were shaken intermittently for 24 h at 65 8C, centrifuged, and the supernatant analysed for cyclodextrin. The desorption procedure was completed three times for each sample.Instrumental and analytical X-Ray diffraction patterns were obtained using a Scintag XDS 2000TM diffractometer, with Cu-Ka radiation (40 kV and 30 mV) and a scanning rate of 18 min21. Sample slides were prepared by pipetting the CMCD–Mg/Al LDH complex suspensions onto glass slides, which were dried at 40 8C for 24 h before analysis. Infrared spectra were obtained by using the KBr disc (1–2% w/w) method and a Michelson Series FT-IR spectrometer with 2 cm21 resolution and 35 scans.Results Isotherm sorption The sorption isotherms of carboxymethyl-b-cyclodextrin by Mg/Al LDH are shown in Fig. 1, which indicates Mg/Al LDH has a stronger affinity for CMCD(3) than CMCD(14). The average maximum sorption capacity (calculated from the sorption plateau) for CMCD(3) was 415 ± 24 mmol kg21 (n = 4), and 222 ± 13 mmol kg21 (n = 5) for CMCD(14), indicating that Mg/Al LDH could sorb approximately twice as much CMCD(3) as compared to CMCD(14). Desorption experiments indicated that 12.1± 5.5% (n = 4) of the CMCD(14) retained by Mg/Al LDH was desorbed at the lower concentration range, while no desorption was observed for samples with maximum sorption.This, to some extent, suggests there was intercalation of CMCD(14) into the Mg/Al LDH interlayers. Similar desorption tendencies were also found for CMCD(3). X-Ray diVraction The X-ray diffraction patterns of the original Mg/Al LDH, its complexes with CMCD(14) (obtained at 3.57 mM solution at 65 8C) and with CMCD(3) (obtained at 12.2 mM solution at 65 8C) are shown in Fig. 2. The original Mg/Al LDH exhibited XRD patterns similar to those reported previously,25,26 with an interlayer spacing (d003) of 8.74 Å; however, Meyn et al.1 reported a basal spacing of 7.4 Å. In addition, the original Mg/ Al LDH contained a small amount of particles with lower interlayer spacings of 7.57 Å [Fig. 2(a)]. The Mg/Al LDH had greater basal spacings, as compared with the values observed for halogen (Cl2 = 7.85, Br2 = 7.95 Å) and carbonatecontaining (7.65 Å) LDHs.26,27 Presumably, the monovalent NO3 2 occupies a space corresponding to three oxygens and is greater than that occupied by other anions such as Cl2 and Br2.26 Nitrate was apparently forced to adopt an arrangement that favoured the closest possible packing. This can lead to strong repulsion inside the interlayer region when the concentration of nitrate increases.Consequently, the LDH exchanged with nitrate had a higher basal spacing value,26 which would presumably favour the CMCD incorporation. The original Mg/ Al LDH also had a peak at 4.82 Å, which was interpreted as an indication of some degree of cation ordering in Mg/Al LDH.28 The XRD results indicated that, for both CMCDs (3 and 14), little d-spacing change was observed until retention reached the sorption plateau, suggesting they could only be intercalated into Mg/Al LDH interlayers at higher concentrations.At higher CMCD loadings, the Mg/Al LDH basal spacing (d003) expanded from 8.74 to 15.48 Å for CMCD(14), and to 20.63 Å for CMCD(3). FT-IR spectra The FT-IR spectra of pure CMCD(3), Mg/Al LDH, and CMCD(3)–Mg/Al LDH complex are shown in Fig. 3. Because CMCD(14) and its complex with Mg/Al LDH exhibited the same patterns, they are not shown here. The assignment of major bands of CMCD(3) [Fig. 3(b)] was made by reference to the study on b-cyclodextrin by Wiedenhof et al.29 and are as Fig. 1 Adsorption of (a) CMCD(14) and (b) CMCD(3) by Mg/Al LDH at 65 8C. Note that the scales are differentJ. Chem. Soc., Dalton Trans., 1997, Pages 1961–1965 1963 follows: 3380 (OH stretching), 2930 (CH stretching), 1605 [ionized carboxyl group, nasym(CO)], 1419 (CH deformation) and 1163–940 (CO/CC stretching). The FT-IR spectrum of the Mg/Al LDH displayed the characteristics bands of LDHs and was analogous to that reported by Miyata.26 The absorption at 3500–3600 cm21 is attributed to the H-bond stretching vibrations of the OH group in the brucite-like layer.In the 600– 1800 cm21 region, there were some bands related to vibrations of the anions and to cation-oxygen vibrations. The major absorption band of NO3 2 appeared at 1384 cm21. Miyata 26 has Fig. 2 X-Ray diffraction patterns of (a) Mg/Al LDH, (b) CMCD(14)– Mg/Al LDH complex, and (c) CMCD(3)–Mg/Al LDH complex Fig. 3 The FT-IR spectra of (a) Mg/Al LDH, (b) CMCD(3) and (c) CMCD(3)–Mg/Al LDH complex reported band splitting in some LDHs containing different anions and suggested that the NO3 2 could exist in the interlayer region of Mg/Al LDH as a mono- or a bi-dentate complex.After the Mg/Al LDH was exposed to CMCD(3), its FT-IR spectrum [Fig. 3(c)] clearly demonstrated that CMCD(3) had been adsorbed by Mg/Al LDH. The principal change occurred with the major peak of NO3 2 (1384 cm21), which decreased sharply due to its exchange by CMCD(3), while the band of ionized carboxyl group (1605 cm21) significantly increased.This indicates that CMCD(3) was adsorbed by Mg/Al LDH through an ion-exchange reaction and that the carboxyl group of CMCD(3) was held through electrostatic attraction between the negative charge of the carboxyl group and the positive charge of the Mg/Al LDH layers. The same conclusions also hold for CMCD(14). Discussion Previous studies 1,9 have indicated that LDHs react easily with various types of organic anions.However, because hydrotalcitetype clays have considerably higher charge densities than silicate-based clays, it was expected that hydrotalcites would be more difficult to swell and for exchange processes to occur. Earlier work30 has shown that the anion-exchange reactions of hydrotalcite proceed with greater difficulty as the charge density of the material increases (i.e., as the Mg: Al atomic ratio decreases) due to increased electrostatic attraction. This was confirmed in our preliminary studies that indicated Ca/Al LDH, which has a higher charge density than Mg/Al LDH,1 had little affinity for CMCD and could not intercalate CMCD molecules.On the contrary, Mg/Al LDH was capable of intercalating CMCD molecules at higher loadings, while the uptake of CMCD at lower concentrations might be contributed to sorption onto the external surfaces of Mg/Al LDH. Cyclodextrin should be regarded as a truncated cone rather than a cylinder.13 There are seven primary and 14 secondary hydroxyl groups along the b-cyclodextrin cavity. On the side where the secondary hydroxyl groups are situated, the diameter of the cavity is larger than on the side with the primary hydroxyls, since free rotation of the latter around the C5]C6 bond reduces the effective diameter of the cavity.b-Cyclodextrin has an approximate torus thickness of 7.8, an outer diameter of 15.3 and an inner diameter of 7.8 Å (as shown in Fig. 4). Primary hydroxyl groups are more readily substituted than are the secondary hydroxyl groups, and the secondary hydroxyls at C3 are much less reactive than those at C2.13 Therefore, for CMCD(3), the three carboxymethyl groups were assumed to be located on the narrow side of the cyclodextrin ring, while for CMCD(14), both primary and secondary hydroxyls at C2 have been substituted by carboxymethyl groups.In deducing the geometry of CMCD between Mg/Al LDH interlayers, we first considered the brucite layer thickness of 4.78 Å,31,32 which left a spacing occupied by CMCD of 10.7 Å for CMCD(14) and of 15.8 Å for CMCD(3).Considering the dimensions of the b-cyclodextrin molecules, the CMCD(14) molecules could only adopt a monolayered arrangement [Fig. 4(a)] with its cavity axes perpendicular to the LDH layer and carboxymethyl groups on both sides of the cyclodextrin cavity attached to the Mg/Al LDH surfaces. As the equivalent area (area per charge) 1 for ideal Mg/Al LDH is 32.4 Å2, the theoretical internal surface area of one face can be estimated 33 to be [(6.02 × 1023) × 32.4]/(333.9 × 1020) = 584 m2 per g LDH {333.9 is equal to the molecular mass of [Mg3Al(OH)8]NO3?2H2O}. Assuming that cyclodextrin molecules are hexagonally closepacked in the monolayer phase, the effective area per molecules was estimated 17 to be 203 Å2, therefore, CMCD(14) content is (584 × 1020)/203[(6.02 × 1023) × 103] = 478 mmol per kg LDH, which is nearly twice the observed value of 222 ± 13 mmol per kg LDH.Alternatively, based on charge balance, we could also obtain a value of 428 mmol per kg LDH by assuming that all of1964 J. Chem. Soc., Dalton Trans., 1997, Pages 1961–1965 the charge sites are occupied by carboxymethyl groups. Therefore, it would be reasonable to assume that the CMCD(14) molecules were loosely packed and existed as discrete isolated entities within the LDH interlayers, calculated to be approximately 6.5 Å apart from each other. This could be due to the conformation of carboxymethyl groups.As none of the cyclodextrin hydroxyls points into the cavity and the C6]O6 bonds are preferentially directed away from the centre of the ring,13,34–36 then the carboxymethyl groups could orient towards the exchange sites outside the cavity, leaving the exchange sites above or below the cavities inaccessible for carboxymethyl groups. This conformation would make the effective area of cyclodextrin larger than 203 Å2.Hence, the larger effective area of cyclodextrin and the competition for exchange sites, as well as the steric hindrance effect, would result in the loose packing of CMCD(14) molecules. By subtracting the surface area occupied by cyclodextrin, we obtained a fairly close value to the observed data, confirming the above reasoning. The thickness of this packing model is 2.9 Å greater than the cyclodextrin torus thickness, which might be ascribed to the carboxymethyl groups along the sides of the cavity.The contribution of carboxymethyl groups on both sides of the cavity was estimated to be 5.5 Å1 (observed value is 2.9 Å). The difference of 2.6 Å might be attributed to the flexibility or the outwards conformation of the carboxymethyl groups. In regard to the CMCD(3) molecules, the space occupied by CMCD(3) is 15.8 Å, which is about twice the torus thickness or equivalent to its outer diameter. Therefore, two types of arrangements are possible: a parallel-monolayered arrangement [Fig. 4(b)] and a perpendicular-bilayer-like coverage [Fig. 4(c)]. Using the same evaluation method described above, the CMCDC(3) content for both the bilayer phase and parallelmonolayer phase amounts to 940 and 813 mmol per kg LDH, Fig. 4 Schematic representation (not to scale) of the possible arrangements for (a) a perpendicular monolayer CMCD(14)–Mg/Al LDH complex, (b) a parallel-monolayer CMCD(3)–Mg/Al LDH complex, and (c) a perpendicular-bilayer CMCD(3)–Mg/Al LDH complex respectively, which was nearly twice as much as the observed data (415 ± 24 mmol per kg LDH).This suggests that CMCD(3) molecules are also loosely packed, owing to the reason described above. However, it is difficult to predict the relative spatial arrangement among CMCD(3) molecules due to the uncertainty of the position and the flexibility of the carboxymethyl groups on the ring side. Based on charge balance, the estimated CMCD(3) content could be as high as 1996 mmol per kg LDH, which is about four times the observed value, indicating that nearly 80% of the charge sites were unoccupied by carboxymethyl groups. This could also be attributed to the blocking of exchange sites by cyclodextrin, as well as the low substitution degree of CMCD(3).In the parallel-monolayer phase, the three carboxymethyl groups could presumably be attached to both LDH layers. In a bilayerlike model, because the CMCD(3) molecules would associate face-to-face with hydrogen bonds between their secondary hydroxyl groups, as observed in several cyclodextrin inclusion compounds,17,36 the contribution of carboxymethyl groups together with the intermolecular hydrogen bonds that have a length 17 of 2.8 Å would make the thickness of the bilayer more than twice that of the torus thickness, or larger than the observed 15.8 Å.Thus, we suggest the parallel-monolayered model is more likely to occur than the bilayer-like model, although a randomly staggered bilayer phase also seems plausible. A definitive explanation, however, awaits further work.Conclusion The following conclusions are apparent from this study. First, carboxymethyl-b-cyclodextrin can be adsorbed by Mg/Al LDH with the retention of CMCD(3) approximately twice that of CMCD(14). Secondly, Mg/Al LDH was able to intercalate CMCD molecules with the intercalated CMCD(14) molecules adopting a monolayer arrangement, with the axis of the cavity perpendicular to the LDH surfaces.Thirdly, Mg/Al LDH and CMCD(3) formed a complex that contained either a perpendicular-bilayer or a parallel-monolayer coverage within the interlayers. Fourthly, both CMCD(14) and CMCD(3) molecules were believed to be loosely packed within the interlayers. Finally, the novel CMCD–Mg/Al LDH intercalates may be of potential use in many aspects. Studies are now in progress to examine the possible use of CMCD–Mg/Al LDH complexes in removal of organic pollutants and chromatographic applications.Acknowledgements This research was supported in part by a student research grant from the Clay Minerals Society (to H. Z.) and in part by the U.S. Air Force grant (93-NA-187). We wish to thank Dr. W. F. Jaynes for his helpful comments on an earlier version of the manuscript, as well as Dr. Keith Carron, University of Wyoming, Chemistry Department for the FT-IR analysis. The valuable suggestions of two anonymous reviewers are also greatly appreciated. References 1 M.Meyn, K. Beneke and G. Lagaly, Inorg. Chem., 1990, 29, 5201. 2 A. Clearfield, M. Kieke, J. Kwan, J. L. Colon and R.-C. Wang, J. Inclusion Phenom. Mol. Recognit. Chem., 1991, 11, 361. 3 K. Chibwe and T. J. Pinnavaia, J. Chem. Soc., Chem. Commun., 1993, 278. 4 K. A. Carrado, J. E. Forman, R. E. Botto and R. E. Winans, Chem. Mater., 1993, 5, 472. 5 T. Kwon, G. A. Tsigdinos and T. J. Pinnavaia, J. Am. Chem. Soc., 1988, 110, 3653. 6 M. A. Drezdzon, Inorg. Chem., 1988, 27, 4628. 7 K. Chibwe and W. Jones, Chem. Mater., 1989, 1, 489. 8 T. Tatsumi, K. Yamamoto, H. Tajima and H.-O. Tominaga, Chem. Lett., 1992, 815.J. Chem. Soc., Dalton Trans., 1997, Pages 1961–1965 1965 9 S. Miyata and T. Kumura, Chem. Lett., 1973, 843. 10 T. Sato and A. Okuwaki, Solid State Ionics, 1991, 45, 43. 11 H. Tagaya, S. Sato, H. Morioka, H. Kadokawa, M. Karasu and K. Chiba, Chem. Mater., 1993, 5, 1431. 12 C. O. Oriakhi, I.V. Farr and M. M. Lerner, J. Mater. Chem., 1996, 6, 103. 13 J. Szejtli, Cyclodextrins and Their Inclusion Complexes, Akademiai Kiado, Budapest, 1982. 14 J. Szejtli, J. Inclusion Phenom. Mol. Recognit. Chem., 1992, 14, 25. 15 M. L. Bender and M. Komiyama, Cyclodextrin Chemistry, Springer, Berlin, 1978. 16 T. Kijima, M. Kobayashi and Y. Matsui, J. Inclusion Phenom., 1984, 2, 807. 17 T. Kijima, J. Tanaka, M. Goto and Y. Matsui, Nature (London), 1984, 310, 45. 18 T. Kijima, J. Inclusion Phenom., 1986, 4, 333. 19 T. Kijima, S. Takenouchi and Y. Matsui, J. Inclusion Phenom., 1987, 5, 469. 20 T. Kijima and Y. Matsui, Nature (London), 1986, 322, 533. 21 T. Kijima, J. Chem. Soc., Dalton Trans., 1990, 425. 22 T. Kijima and K. Ohe, J. Chem. Soc., Dalton Trans., 1992, 2877. 23 P. Mondik, A. Sopkova, G. Suchar and T. Wadsten, J. Inclusion Phenom. Mol. Recognit. Chem., 1992, 13, 109. 24 R. Allmanmn, Chimia, 1970, 24, 99. 25 F. Cavani, F. Trifiro and A. Vaccari, Catal. Today, 1991, 11, 173. 26 S. Miyata, Clays Clay Miner., 1975, 23, 369. 27 S. Miyata, Clays Clay Miner., 1983, 31, 305. 28 M. C. Gastuche, G. Brown and M. M. Mortland, Clay Miner., 1967, 7, 177. 29 N. Wiedenhof, J. N. J. J. Lammers, C. L. Van Panthaleon and B. Van Eck, Die Stärke, 1969, 21, 119. 30 S. Kikkawa and M. Koizumi, Mater. Res. Bull., 1982, 17, 191. 31 Nipponkagakukai, Kagakubinran Kisohen. 2, Maruzen, Tokyo, 1966, vol. 2, p. 1265. 32 R. W. G. Wykoff, Crystal Structure, John Wiley and Sons, New York, 1963, vol. 1, p. 268. 33 H. Kopka, K. Beneke and G. Lagaly, J. Colloid Interface Sci., 1988, 123, 427. 34 W. Saenger, in Inclusion Compounds, eds. J. L. Atwood, J. E. D. Davies and D. D. MacNicol, Academic Press, New York, London, 1984, vol. 2, ch. 8, pp. 231–258. 35 L. F. Lindoy, The Chemistry of Macrocyclic Ligand Complexes, Cambridge University Press, Cambridge, 1989, p. 164. 36 W. Saenger, Angew. Chem., 1980, 19, 344. Received 8th November 1996; Paper 6/07604A

ISSN:1477-9226    

DOI:10.1039/a607604a

出版商:RSC    

年代:1997

数据来源: RSC

摘要:

DALTON J. Chem. Soc., Dalton Trans., 1997, Pages 1967–1972 1967 Complexation of aluminium(III) with several bi- and tri-dentate amino acids Tamás Kiss,*,a Imre Sóvágó,b Imre Tóth,b Andrea Lakatos,a Roberta Bertani,c Andrea Tapparo,d Giorgio Bombi d and R. Bruce Martin e a Department of Inorganic and Analytical Chemistry, Jozsef Attila University, H-6701, Szeged, PO Box 440, Hungary b Department of Inorganic and Analytical Chemistry, Kossuth University, H-4010 Debrecen, PO Box 21, Hungary c Centro di Studio sulla Chimica e Tecnologia dei Compositi Metallorganici degli Elementi di Transizione, CNR, Via Marzolo 9, I-35131 Padova, Italy d Dipartimento di Chimica Inorganica Metallorganica Analitica, Via Marzolo 1, I-35131 Padova, Italy e Department of Chemistry, University of Virginia, Charlottesville, VA 22903, USA Complex formation between AlIII and various bidentate and potentially tridentate amino acids (Gly, Ser, Thr, Gln, Asn, Glu and Asp) and some model compounds (N-acetylaspartic acid, succinic acid, and 2-sulfanylsuccinic acid) was studied in aqueous solution by pH-potentiometric and multinuclear (1H, 13C and 27Al) NMR techniques.A weak interaction through the carboxylate function was unambiguously detected for all of these ligands. For Asp, however, which contains two carboxylates and a central amino donor, tridentate co-ordination is strongly suggested. Although its steric environment is similar to that of the amino function in Asp, binding of the thiolate sulfur of 2-sulfanylsuccinic acid could not be verified.Aluminium is known to form fairly stable complexes with negatively charged O-donor chelating biomolecules such as aliphatic and aromatic hydroxycarboxylates, catecholates, phosphates, etc. Its binding ability to the protein building block a-amino acids, however, has not yet been established unambiguously.1–5 The a-carboxylate group of amino acids is weakly basic (pK ª 2.2), which suggests a rather weak aluminium(III) binding capability.The usual sample composition in the pH-potentiometric titration method (a metal-ion concentration of a few mmol dm23 with a metal ion to ligand ratio of from 1 : 1 to 1 : 5) is not suitable for the detection of complex formation.1 It has been found that the strength of complexation to AlIII decreases strongly in the sequence dicarboxylic acid @ hydroxycarboxylic acid > carboxylic acid @ amino acid. The weakening effect of amino substitution was explained in terms of the electrostatic repulsive effect of the NH3 1 group.1 Use of an LFER (linear free-energy relation) approach led to a log b value of 5.8 being estimated for Al31 and glycinate.3 According to this estimate, co-ordination/chelation of Al31 should become distinguishable from hydrolysis at glycine (Gly) concentrations greater than 25 mmol dm23.Precise pH-metric studies have been carried out by Dayde 4 in several AlIII–amino acid (Gly, Ser, Thr, Asp, Glu and His) systems at 37 8C and I = 0.15 mol dm23 (NaCl) and at a high ligand concentration of 0.02 mol dm23.Stability constants for 1 : 1 complexes were obtained by parallel refinement of the stability constants of the AlIII–hydroxo and –amino acid complexes. A log KAlA value in the range 5.7–6.2 was determined for the bidentate amino acids. In contrast, Yadava and co-workers,2 detected the formation of much stronger AlA, AlA2 and AlA3 complexes with Leu, Ser and Thr by using an ionophoretic technique, although admitted the rather limited precision of the method.Rao and Rao5 recently made a qualitative characterization of the interaction between Al31 and Ala. With potentially tridentate amino dicarboxylic acids such as Asp and Glu a somewhat stronger interaction could be detected: the formation of various protonated species, and ternary dinuclear hydroxo species, was assumed.4–6 In the present paper we evaluate the stability constants for the binding of AlIII to various amino acids, such as Gly, Ser, Thr, Gln, Asn, Glu and Asp.A comparative study was carried out with N-acetylaspartic acid (Ac-Asp), succinic acid (H2succ) and 2-sulfanylsuccinic acid (H3Ssucc). The speciation results of the detailed pH-potentiometric measurements were confirmed by 1H, 13C and 27Al NMR measurements. Experimental Reagents The amino acids, the best quality compounds available from Reanal Co., were purified by recrystallization.The compounds H2succ and H3Ssucc were Aldrich products. Their purities were checked, and when possible the exact concentrations of their solutions (prepared freshly each day) were determined by the Gran method.7 The stock solution of Al31 was prepared from recrystallized AlCl3?6H2O and its metal concentration was determined gravimetrically via its quinolin-8-olate. To prevent hydrolysis, the metal-ion stock solution contained 0.1 mol dm23 HCl. The ionic strength of all solutions measured was adjusted to 0.20 mol dm23 KCl.In all cases the temperature was 25.0 ± 0.1 8C. pH-metric measurements The stability constants of the proton and aluminium(III) complexes were determined by pH-metric titration of 25.0 cm3 samples. Owing to the weak complexation reactions, measurements were performed at high total concentrations and at high excesses of the amino acids, 0.02 or 0.04 mol dm23, and the metal ion to amino acid ratio was 1 : 10, 1 : 15, 1 : 20, 1 : 25, 1 : 30 or 1 : 40.Owing to solubility limitations, Asp was measured at 0.016 mol dm23. The titrations were performed until precipitation occurred, with carbonate-free KOH solutions of known concentration (ca. 0.2 mol dm23) under a purified argon atmosphere. Depending on the amino acid and the metal ion to1968 J. Chem. Soc., Dalton Trans., 1997, Pages 1967–1972 amino acid ratio, precipitation occurred in the range pH 4.5– 5.8. The reproducibility of the titration curves was within 0.010 pH units throughout the whole pH range.When equilibration could not be reached in 10 min, titration points were omitted from the calculations. The pH was measured with a Radiometer PHM 84 instrument with a GK2322C combined glass electrode, which was calibrated for hydrogen-ion concentration according to Irving et al.8 The ionic product of water was found to be pKw = 13.76. The concentration stability constants bpqr = [MpAqHr]/[M]p[A]q[H]r were calculated with the aid of the computer program PSEQUAD.9 For the hydroxo complexes of AlIII, the stability constants (log b) assumed10 were 25.52 for [AlH21]21, 213.57 for [Al3H24]51, 2109.1 for [Al13H232]71 and 223.46 for [AlH24]2.NMR measurements Proton, 13C and 27Al NMR spectra were recorded at 25 8C on Bruker WP 200SY and AC200 spectrometers in D2O. The pD values were calculated via the equation pD = pH-meter reading 1 0.4.11 Quantitative 27Al NMR spectra were recorded using the ‘absolute integration mode’, in order to allow comparison of the integral values obtained for the different samples. Typical NMR parameters were: pulse width = 5 ms (308), pulse repetition time = 1.262 s.Results and Discussion Interactions of AlIII with bidentate amino acids (Gly, Ser, Thr, Asn and Gln) Speciation results. Owing to the expected rather weak interactions of AlIII with bidentate aminocarboxylates and its high hydrolytic tendency, very high amino acid excesses were applied in the titration samples.Titration curves of the acid alone, Al31 alone and the Al31–Gly sample at a 1 : 15 metal ion to acid ratio are depicted in Fig. 1. These curves indicate that the pH range of metal–ligand complexation is well separated from the deprotonation (buffering) pH range of the amino acid, and thus the buffering effect of the large excess of the latter is negligible in this pH range (see curves 1 and 2). At the same time the overlap with the hydrolytic processes of AlIII is considerable (see curves 1 and 3).The small, but consequent difference between curves 2 and 3, however, can be unambiguously ascribed to the proton-displacement reactions of AlIII, i.e. to the reaction Al31 1 HA AlA 1 H1. Accordingly, the metal–ligand formation constants can be determined with acceptable precision.12 Titration curves obtained at high amino acid excess were evaluated by assuming different equilibrium models. The best fit between the experimental and the calculated titration curves was obtained with the species and stability constants listed in Table 1.The pH ranges included in the calculation are also given. It is seen that these AlIII–amino acid systems can be fairly well described by the same speciation model, assuming a 1 : 1 complex [AlA]21, its deprotonated form [AlAH21]1, and the dinuclear species [Al2AH21]41. Other complexes, such as [Al(HA)]31, and all 1 : 2 complexes, were rejected by the computer program.As an illustration, species distribution curves for the AlIII–Gly system are depicted in Fig. 2. Proton, 13C and 27Al NMR measurements. Although the pH-metric results are fairly convincing, independent 1H, 13C and 27Al NMR measurements were also made unambiguously to demonstrate this weak interaction between AlIII and simple bidentate amino acids. First, the 1H and 13C NMR spectra of Fig. 1 Titration curves of samples containing Al31 and/or Gly: 1 (s), cGly = 0.04 mol dm23; 2 (1), cGly = 0.04, cAl = 0.0027 mol dm23; 3 (d), cAl = 0.0027 mol dm23 Fig. 2 Species distribution curves in the AlIII–Gly system. cGly = 0.16, cAl = 0.008 mol dm23 Table 1 Proton (log K) and aluminium(III) complex-formation constants (log b) of several bidentate amino acids at 25 ± 0.01 8C and I = 0.20 mol dm23 (KCl) log K(NH2) log K(CO2 2) [AlA]21 [AlAH21]1 [Al2AH21]41 Fitting * No. of points pH Range Al 1 HA AlA 1 H AlA AlAH21 1 H Gly 9.57(1) 2.34(2) 5.91(10) 1.08(9) 4.35(9) 0.0108 248 2.6–4.9 23.66 24.83 Ser 9.02(1) 2.16(2) 5.66(11) 0.62(23) 3.75(11) 0.0109 160 2.3–5.4 23.36 25.04 Thr 8.91(1) 2.15(2) 5.51(12) 0.94(15) — 0.0130 80 2.6–4.8 23.40 24.57 Gln 8.99(1) 2.16(2) 5.61(13) 1.33(4) — 0.0058 84 2.3–4.7 23.38 24.28 Asn 8.71(1) 2.14(2) 5.50(28) 1.31(8) — 0.0134 168 2.3–5.2 23.21 24.19 * Average difference in the calculated and experimental titration curves expressed in cm3 of the titrant.J. Chem.Soc., Dalton Trans., 1997, Pages 1967–1972 1969 the H–Gly and AlIII–Gly systems were taken at high concentrations (cAl = cGly = 0.5 mol dm23) at pH ª2.5 (the highest pH before precipitation occurred).The spectra were found to consist of two clearly separated sets of signals: the complexation resulted in a well detectable 0.07 ppm downfield shift in the 1H(CH2) signal [see Fig. 3(a)], while a 0.46 ppm downfield shift in the 13C(CH2) signal and a 1.0 ppm upfield shift in the 13C(CO) signals (see Fig. 4).As is seen in Fig. 3(a), under these conditions nearly 50% of the total Gly is bound to AlIII. With dilution, of course, dissociation of the complex increases, but as shown in Fig. 3(b) complex formation can be unambiguously detected even in 0.05 mol dm23 solutions. The separated signals indicate a ‘slow exchange’ range for the complex and the free amino acid system. Substantial broadening of the signals of the free amino acid is observed, especially at pH 3.5 for the more dilute sample.This cannot be explained by a mutual two-site exchange between the complex and free amino acid, as the equality p1(Dn2� 1 )1 = p2(Dn2� 1 )2 (where p is the molar ratio of each site and Dn2� 1 is the half width of the original) is not fulfilled. A detailed study of the kinetics of these processes is under progress. At mmol dm23 aluminium(III) concentrations high Fig. 3 Proton NMR spectra of the Al31–Gly system at 2.5 (a) and 3.5 pH (b); cAl = 0.5, cGly = 0.5 mol dm23 Fig. 4 Carbon-13 NMR spectra of the Al31–Gly system; cAl = 0.5, cGly = 0.5 mol dm23, pH 2.44 excesses of amino acid have to be applied to enhance the AlIII– amino acid interaction, and thus complex formation cannot be monitored using the ligand nuclei; accordingly 27Al NMR measurements were carried out. Since AlIII is quadrupolar, asymmetry in the ligand field produces an increase in the NMR linewidth. The resonances of low-symmetry aluminium(III) species can be so broad (linewidth higher than a few thousand Hz) that they cannot be detected in practice as they merge into the baseline.When the symmetry is high, e.g. in [Al(H2O)6]31, the 27Al spectral peak is sharp: in an 8 × 1023 mol dm3 Al31 solution at pH 1.26 a single peak with a linewidth of 4 Hz was measured. For the same solution at pH 4.08 a single peak still occurred at d 0.71, but with a much larger linewidth, while at pH 4.51 three peaks could be detected: at d 1.14 (Dn2� 1 = 100), at 63.6 (12) and a very broad band (Dn2� 1 > 1000 Hz) centred around d 30.The peak at d ª1 relates to AlIII in a fairly symmetrical octahedral environment; the other sharp signal, at around d 63.6, can be ascribed to the central tetrahedral AlIII of the tridecanuclear hydroxo species,13 while the remainder of the AlIII, in a much less symmetrical environment, is reflected by the broad band. Area integration revealed that in this sample ª60% of the Al is involved in peak 1, ª10% in peak 3 and the rest in peak 2.In a similar way, we used area integration to estimate the amount of AlIII found in the different chemical environments in solutions containing a high excess of amino acid (Gly or Ser) besides AlIII. The results listed in Table 2 show that at pH ª3.0 the presence of a high excess of amino acids has practically no effect on the 27Al NMR spectrum: all the AlIII is in octahedral species {mostly [Al(H2O)6]31}. At pH 4.5, however, in the presence of a 20-fold excess of the amino acids, 20% less AlIII is to be found in octahedral environment. At pH ª5 the solution is still clear [whereas in the absence of ligands Al(OH)3 precipitates from the solution], and only AlIII bound in the tridecanuclear species can be clearly detected, but most of the AlIII Table 2 Aluminium-27 NMR spectral data for AlIII–amino acid interaction cAl/mol dm23 pH d Dn2� 1 /Hz Area Al31 solution 0.008 0.008 0.008 1.26 4.08 4.51 0.50 0.71 1.14 63.6 4 50 100 12 100 104 60 10 AlIII–Gly 1 : 20 ratio 0.008 0.008 0.008 0.008 3.09 4.08 4.49 5.02 0.55 0.96 1.20 63.6 63.5 6 79 120 24 40 100 98 40 24 AlIII–Ser 1 : 20 ratio 0.008 0.008 0.008 0.008 3.01 3.96 4.47 5.00 0.50 0.86 0.93 63.7 63.7 5 73 144 26 35 110 100 41 34 AlIII–Asp 1 : 5 ratio 0.003 0.003 0.003 4.08 4.47 5.02 0.93 9.9 1.37 9.9 10.0 90 130 120 130 180 36 29 12 29 14 AlIII–H3Ssucc 1 : 5 ratio 0.003 0.003 0.003 3.53 4.02 4.53 0.73 0.90 1.70 44 110 140 60 25 Precipitate1970 J.Chem. Soc., Dalton Trans., 1997, Pages 1967–1972 must be in a less symmetrical environment and this gives a very broad, hardly detectable band. Comparison of the speciation curves (Fig. 2) and the 27Al NMR data (Table 2) indicates a reasonably good correlation between the two. At pH ª3.0, where AlIII is almost exclusively in the free form, the total aluminium yields a signal at around d ª0.5. The somewhat larger linewidth as compared with that of [Al(H2O)6]31 under more acidic conditions can be explained by the effect of the relatively slow ligand-exchange reaction.A similar broadening of the signals was observed by Garrison and Crumbliss 14 for the AlIII–hydroxamic acid systems. When the pH is increased to ª4 or ª4.5 increasing amounts of AlIII disappear from the octahedral aluminium(III) band at d ª1 as compared with the intensity-distribution data obtained for the ligand-free Al31 solution. Similarly, the relative amount of the tridecanuclear hydroxo species also decreases. The ‘missing’ AlIII is probably bound to amino acid in a fairly unsymmetrical geometry, thereby resulting in a very broad, hardly detectable 27Al NMR band. At pH ª5.0 only the tridecanuclear species gives a sharp signal in the presence of amino acids.As its relative intensity is somewhat higher than at pH 4.5, OH2 tends to displace simple amino acids from the co-ordination sphere of AlIII with increasing pH. The speciation and the multinuclear NMR data unambiguously demonstrate the interaction between AlIII and simple amino acids under weakly acidic conditions at the mmol dm23 aluminium(III can be said about the bonding mode in the aluminium( III) complexes? As the affinity of AlIII for N-donors is very low,3 it is not too likely that the NH2 group of these amino acids is involved in aluminium(III) binding. Hence, it is probably in protonated form (NH3 1) in weakly acidic solutions. Accordingly, [AlA]21, the first complex formed, must be a mixed-ligand hydroxo complex [Al(HA)(OH)]21. The other microform [AlA]21 containing a five-membered chelate may be a minor form.The formation constants obtained for the species are in fairly good agreement with those determined by Dayde.4 The species [AlA]21 loses a proton with pK ª4.2–5.0, which should be ascribed either to the ionization of a second water molecule in the co-ordination sphere of AlIII or to deprotonation and coordination of the amino group.The complex, however, is only a minor species. The fairly low pKAlA values {0.5–1.0 log unit lower than pK for [Al(H2O)6]31 = 5.52; see Table 1}, i.e. the slight increase in acidity of the water molecules in the complex should indicate either (i) involvement of the amino group in the co-ordination (see above), or (ii) a change in the co-ordination geometry from octahedral to tetrahedral during the water ionization process. It must be mentioned that a similar acidification of the co-ordinated water molecules was observed in the case of AlIII–adenosine 59-phosphate complexes.15 It is seen in Fig. 2 that the mononuclear species [AlA]21 is formed in parallel with a dinuclear complex [Al2AH21]41 in the same pH range. It was found that these two species could partly replace each other in the calculation. The fitting parameter changed only a little, when besides [AlAH21]1, only one of them was assumed in the calculation. The stability data listed in Table 1 were obtained when both species were simultaneously included in the speciation model.The ion [Al2AH21]41 is very probably a carboxylate- and dihydroxo-bridged complex with a protonated ammonium group, i.e. [Al2(OH)2(HA)]41. A complex with this binding mode has also been observed in AlIII–monocarboxylic acid 16 and –carbonate systems.17 As this structure is found in minerals containing dihydroxo-bridged aluminium chains, Öhman16 assumes that it is a generally occurring structure for all dinuclear AlIII–monocarboxylato complexes.He demonstrated that the ability of AlIII to form this complex increases with increasing pKCO2H and established approximate linearity between the log K values characterizing the formation reaction 2Al31 1 A Al2AH22 1 2H1 and the pKCO2H values. The above tendency exists for the aminocarboxylates Gly and Ser, as the significantly lower basicity of the carboxylate functions of the amino acids (pK ª2.1–2.3) results in a lower stability of the dinuclear complexes.Aluminium(III) interactions with Glu, Asp and some related compounds Speciation results. Glutamic acid and aspartic acid, containing carboxylic functions (which have a high affinity for AlIII) at both ends of the molecule, can be potentially more efficient binders of AlIII than the amino acids discussed so far. With transition-metal ions, where the basic bonding mode of aamino acids is five-membered (CO2 2, NH2) chelation, Asp readily acts as a tridentate ligand, with formation of a five- 1 sixmembered (CO2 2, NH2, CO2 2) joint chelate system.The tridentate co-ordinating ability of Glu is much weaker, due to the lower stability of the seven-membered second chelate ring. Thus, Glu rather acts as a bidentate ligand showing similarity with simple amino acids. Aspartic acid (Asp) has been shown to form (CO2 2, CO2 2) chelates (via formation of a sevenmembered ring), especially with hard metal ions.18 Although AlIII has low affinity for N-donors, in Asp, where the amino group is adjacent on both sides to strong AlIII-binder carboxylate functions, metal ion-induced deprotonation of the NH3 1 group and co-ordination of the amino group may be more Table 3 Proton (log K) and aluminium(III) complex-formation constants (log b) of several tridentate amino acids and related compounds at 25 ± 0.01 8C and I = 0.20 mol dm23 (KCl) log K(NH2) log K(CO2 2) log K(CO2 2) [AlHA]21 [AlA]1 [AlAH21] [AlAH22]2 [Al2A]41 Fitting b No.of points pH Range Al 1 HA Al(HA) AlHA AlA 1 H AlA AlAH21 1 H AlAH21 AlAH22 1 H Al 1 H2A(H) AlA(H) 1 2H Al 1 H2A(H) AlA(H)H21 1 3H Glu 9.50(1) 4.09(2) 2.04(3) 10.88(22) 7.29(4) 2.55(3) — 9.46(15) 0.0092 244 2.6–5.0 1.38 23.59 24.74 — 24.75 28.34 Asp 9.62(1) 3.66(1) 1.94(5) 11.76(6) 7.87(4) 3.30(3) 22.32(7) — 0.0128 201 2.5–5.9 2.14 23.89 24.57 25.62 23.46 27.35 Ac-Asp — 4.49(2) 3.04(2) 6.15(3) 2.24(9) 21.91(14) —— 0.0075 158 2.0–4.6 1.66 23.91 24.15 — 25.29 29.44 H2succ — 5.24(2) 3.96(2) 7.03(15) 3.63(8) 20.53(5) 25.55(9) — 0.0086 223 2.0–5.9 1.79 23.40 24.16 25.02 25.57 29.73 H3Ssucc a 10.20(3) 4.47(2) 3.05(2) 12.99(2) 8.63(9) 4.05(4) —— 0.0128 113 2.5–5.0 2.79 24.36 24.58 — 24.73 29.09 a As the neutral form of H3Ssucc is H3A, while that of the other ligands is H2A the charge of the corresponding complexes of H3Ssucc is always one less.b Average difference in the calculated and experimental titration curves expressed in cm3 of the titrant.J.Chem. Soc., Dalton Trans., 1997, Pages 1967–1972 1971 likely, similarly to alcoholic OH and amide NH groups which exhibited an enhanced metal-binding ability when they were in a central position within the molecule.19,20 To obtain information about the binding ability of the NH2 group in Asp, its Nacetyl derivative (with a blocked amino group) and succinic acid (without the central amino group) were also studied. An attempt was first made to evaluate the pH-metric titration data on AlIII–Glu and –Asp systems with the same speciation model as accepted for simple amino acids, taking into account that the ligand contains an extra proton on the terminal carboxylic function. For Glu a reasonably good fit was obtained with the species and stability constants listed in Table 3.However, the AlIII–Asp system gave a very poor fit with this model, especially in the high-pH range (the AlIII–Asp system could be measured up to pH ª6 without precipitation or unstable pH-meter readings, while in the AlIII–Glu system precipitation occurred at pH ª5).When another deprotonated complex, [AlAH22], was also included in the model a reasonable fit was obtained between the experimental and calculated titration data. The aluminium(III) systems of the two reference compounds N-Ac-Asp and H2succ were fitted with the same speciation model as for Asp, with the exception that in the case of the Ac-Asp system the pH range evaluated was about one unit narrower due to precipitation.The results of the computer evaluation of the pH-metric data are shown in Table 3. Species distribution diagrams for the AlIII–H2succ, –Ac-Asp and –Asp systems are depicted in Figs. 5–7. The complex [Al(HA)]21 found for H2succ and Ac-Asp is a monodentate carboxylate-co-ordinated species; its formation constant {Al31 1 AH2 [Al(HA)]21} is in reasonably good agreement with those obtained by Öhman16 for simple aliphatic and aromatic carboxylates.With these ligands, unlike simple amino acids, the absence of the positively charged NH3 1 group will not hinder the formation of such monocarboxylate- co-ordinated species. The consecutive loss of two protons from this [Al(HA)]21 species can be ascribed to the Fig. 5 Species distribution curves in the AlIII–H2succ system; csucc = 0.015, cAl = 0.003 mol dm23 Fig. 6 Species distribution curves in the AlIII–Ac-Asp system; cAc-Asp = 0.015, cAl = 0.003 mol dm23 formation of a seven-membered (CO2 2, CO2 2) chelate and the ionization of a co-ordinated water molecule.The basicityadjusted stability constants, which take into account the differences in basicity of the co-ordinating donor groups, are practically the same for Ac-Asp and H2succ, confirming the same binding modes in their corresponding complexes. The presence of the bulky N-acetyl moiety in Ac-Asp, however, will result in the early formation of a precipitate, which can be either Al(OH)3 or perhaps the neutral [AlAH21] complex, from an oversaturated solution.When the AlIII-binding properties of H2succ and Ac-Asp are compared with those of Asp it can be seen that Asp does not form monodentate CO2 2-co-ordinated species (its stoichiometry should be [AlAH2]31), because of the electrostatic repulsive effect of the protonated NH3 1 group. However, when the stability constants of the (CO2 2, CO2 2) chelates are compared Asp displays significantly higher stability.In the case of Asp this species has a stoichiometry of [Al(HA)]21, due to the protonated NH3 1 group. Assuming only bidentate (CO2 2, CO2 2) co-ordination in the [Al(HA)]21 (with protonated NH3 1) and [AlA]1 (mixed-ligand hydroxo species with protonated NH3 1) complexes of Asp, the corresponding basicity-adjusted stability (proton-displacement) constants are about two orders of magnitude larger for Asp (see the last two rows of Table 3), which strongly suggests the involvement of the NH2 group in binding to AlIII. Hence, we can assume two microforms for both Asp complexes, [Al(HA)]21 and [AlA]1 (see Scheme 1).A similar tridentate (CO2 2, NH2, CO2 2) co-ordination of iminodiacetate to AlIII was demonstrated by multinuclear NMR21,22 and X-ray analysis in the solid state.22 It is worthy of mention that even the corresponding Glu complexes exhibit an enhanced stability, and thus NH2 involvement via the formation of the microform with a five- 1 seven-membered joint chelate Fig. 7 Species distribution curves in the AlIII–Asp system; cAsp = 0.015, cAl = 0.003 mol dm23 Scheme 11972 J. Chem. Soc., Dalton Trans., 1997, Pages 1967–1972 system can also be assumed. The formation of the complex [AlAH22]2 likewise suggests co-ordination of the amino group, as this stoichiometry can be ascribed to only one reasonable binding mode: (CO2 2, NH2, CO2 2) tridentate co-ordination of the ligand and two OH2 groups around the metal ion.Mixed hydroxo complexes of AlIII with three OH2 groups have never been detected with any ligand. Aluminium-27 NMR measurements. The multinuclear NMR measurements point to the above discussed binding properties of the ligands. In contrast with simple amino acids, the 27Al NMR spectra indicate a much stronger interaction (see Fig. 8) with Asp. The relatively sharp signal (Dn ª130–180 Hz) at around d 10 suggests octahedral AlIII in a fairly symmetrical chemical environment in the AlIII–Asp complex formed at pH >4.Succinic acid reveals somewhat similar 27Al NMR; behaviour. At pH 4.2 it displayed two peaks, at d 0.5 and 4.6; the former is due to free Al31, while the latter, broader resonance is attributed to the [AlA]1 with (CO2 2, CO2 2) chelation. From a study of the 27Al NMR spectra of a series of AlIII–aminopolycarboxylate complexes, Karweer et al.21 demonstrated a nearly linear relationship between the chemical shift and the denticity of the ligand.They established that replacement of a H2O molecule by a nitrogen or oxygen donor with the formation of five-membered chelate rings produces an average deshielding of 6–7 ppm, while replacement of H2O by OH2 in octahedral geometry results in deshielding by ª2 ppm. Although no data have been reported for the deshielding effect of the formation of six- or seven-membered chelates, it is reasonable to assume that it decreases with increasing ring size (weaker interaction).Accordingly, the resonance at d 4.6 for seven-membered (CO2 2, CO2 2) co-ordination (AlIII–H2succ system) and that at d 10.0 for five- 1six-membered (CO2 2, NH2, CO2 2) chelation (AlIII–Asp system) conform with the relationship reported by Karweer et al.21 2-Sulfanylsuccinic acid, which is a thiol analogue of Asp where the NH2 group is replaced by an SH group, displays similarity to H2succ in its AlIII-binding properties. The protondisplacement constants listed in the last two rows of Table 3 indicate a slightly enhanced stability of the complexes [AlA] Fig. 8 Aluminium-27 NMR spectra of AlIII–Asp at 1 : 5 metal to amino acid ratio at different pH values; cAl = 0.005 mol dm23 and [AlAH21]2 as compared with these of H2succ. However, this is even less than was found for AlIII–Glu complexes; hence, we assume that (CO2 2, CO2 2) chelation is the basic bonding mode in the AlIII–H3S succ complexes. Conclusion In accordance with earlier reports, simple a-amino acids proved to be weak binders of Al31.However, their complexation becomes distinguishable from hydrolysis of Al31 at high aluminium concentrations (0.5 mol dm23 at a 1 : 1 metal ion to amino acid ratio) or at mmol dm23 Al31 with amino acid concentrations greater than 20 mmol dm23, and it could be detected unambiguously by means of both pH-potentiometric and 1H, 13C and 27Al NMR measurements. With the tridentate Asp, which contains two carboxylate-O and one central amino-N binding donors, a much stronger interaction was observed.No such strong complexation could be detected with either H2succ or Ac-Asp, which lack the central amino binding site. These results strongly suggest that, besides the negatively charged CO2 2 donors, in the event of a favourable steric arrangement, the NH2 group (not only in amino acids, but probably also in peptides and proteins) can likewise participate in the binding of AlIII.The binding ability of thiolate S2, however, could not be verified in H3S succ, although it occurs in a similarly favourable steric environment. Acknowledgements This work was supported by the Hungarian Academy of Sciences (Project No. OTKA T7458/1993 and T14075/1995). References 1 E. Marklund and L. O. Öhman, Acta Chem. Scand., 1990, 44, 353. 2 S. Singh, S. Gupta, P. C. Yadava, R. K. P. Singh and K. L. Yadava, Z. Phys. Chem. Leipzig, 1986, 267, 902. 3 R. B. Martin, in Aluminium in Chemistry, Biology and Medicine, eds. M. Nicolini, P. F. Zatta and B. Corain, Cortina International, Verona, 1991. 4 S. Dayde, Ph.D. Thesis, Université de Toulouse, 1990. 5 K. S. J. Rao and G. V. Rao, Mol. Cell. Biochem., 1994, 137, 61. 6 Ph. Charlet, J. P. Deloume, G. Duc and G. T. David, Bull. Soc. Chim. Fr., 1984, I-222. 7 G. Gran, Acta Chem. Scand., 1950, 29, 559. 8 H. Irving, M. G. Miles and L. D. Pettit, Anal. Chim. Acta, 1967, 38, 475 . 9 L. Zékány and I. Nagypál, in Computational Methods for the Determination of Stability Constants, ed. D. Legett, Plenum, New York, 1985. 10 L. O. Öhman and S. Sjöberg, Acta Chem. Scand., Sect. A, 1982, 36, 47. 11 R. G. Bates, Determination of pH, Wiley, New York, 1964. 12 I. Fábián and I. Nagypál, Talanta, 1982, 29, 71. 13 C. Orvig, in The Coordination Chemistry of Aluminum, ed. G. H. Robinson, VCH, New York, 1993. 14 J. M. Garrison and A. L. Crumbliss, Inorg. Chem., 1987, 26, 3660. 15 T. Kiss, I. Sóvágó and R. B. Martin, Inorg. Chem., 1991, 30, 2130. 16 L. O. Öhman, Acta Chem. Scand., 1991, 45, 258. 17 L. O. Öhman and S. Sjöberg, J. Chem. Soc., Dalton Trans., 1983, 2513. 18 T. Kiss, in Biocoordination Chemistry, ed. K. Burger, Ellis Horwood, New York, 1990. 19 T. Kiss, Cs. Simon and Zs. Vachter, J. Coord. Chem., 1987, 16, 225. 20 I. Sóvágó, in Biocoordination Chemistry, ed. K. Burger, Ellis Horwood, New York, 1990. 21 S. B. Karweer, B. P. Pillai and R. K. Iyer, Magn. Reson. Chem., 1990, 28, 922. 22 S. P. Petrosjanc, M. A. Maljarik and A. B. Iljuhin, Zh. Neorg. Khim., 1995, 40, 769. Received 6th December 1996; Paper 6/08250E

ISSN:1477-9226    

DOI:10.1039/a608250e

出版商:RSC    

年代:1997

数据来源: RSC

摘要:

DALTON J. Chem. Soc., Dalton Trans., 1997, Pages 1973–1979 1973 Synthesis, structural and electronic characterisation of trans- [OsCl2(PEt2Ph)3{(NC)2C] C(CN)OH}], a complex featuring a redox-active, tetracyanoethylene-derived ligand ‡ Alexander J. Blake, Lockhart E. Horsburgh, Martin Schröder*,† and Lesley J. Yellowlees * Department of Chemistry, University of Edinburgh, West Mains Road, Edinburgh EH9 3JJ, UK Reaction of mer-[OsCl3(PEt2Ph)3] with AgBF4 and tetracyanoethylene (tcne) in tetrahydrofuran yielded the unexpected product trans-[OsCl2(PEt2Ph)3(tcva)], where tcva is tricyanovinyl alcohol (1,1,2-tricyano-2- hydroxyethylene). Structural characterisation of trans-[OsCl2(PEt2Ph)3(tcva)] was accomplished through IR spectroscopy, fast atom bombardment mass spectrometry, elemental analysis and single-crystal X-ray diffraction.The crystallographic study revealed that tcva is s-bound to the Os atom via nitrogen and that the OH group occupies the cis position with respect to the metal-bound nitrogen. The frontier orbitals of trans- [OsCl2(PEt2Ph)3(tcva)] were investigated by electrochemistry and UV/VIS, IR and EPR spectroelectrochemical studies, which revealed that the complex is best described as trans-[OsIICl2(PEt2Ph)3(tcva)0]. The co-ordination chemistry of tetracyanoethylene (tcne) and 7,7,8,8-tetracyanoquinodimethane (tcnq) have been the source of much interest in recent years.1 They are potent electron acceptors and their reactions with transition-metal complexes exhibit considerable diversity of structural form and electronic interactions. The products of these reactions may be divided into several categories: (i) salts, in which the tetranitrile is present as a counter anion.2–5 The electron-accepting properties of tcne and tcnq are demonstrated by the existence of species such as [{Co(C5Me5)2}2][tcne],2 in which the tetranitrile is present as a dianion.Ferromagnetism 3,4 and other unusual magnetic and electronic properties 5 have often been observed in such species.(ii) Charge-transfer complexes, in which electron transfer from metal to tetranitrile is incomplete. Perhaps the best known example of such a species is [Fe(C5H5)2?tcne].6 One feature of interest in both salts and charge-transfer complexes of tcne/q is the tendency to form low-dimensional solids, in which the physical properties show an unusually high degree of anisotropy.4,7 (iii) Metal–alkene bound p complexes. This has been the most widely reported form of bonding between tcne and transitionmetal complexes.1,8 As a result of the electron-withdrawing CN groups, there is usually considerable back bonding from the metal to the alkene p* orbital.In general, the planarity of the tetranitrile ligand is lost, and considerable lengthening of the alkene double bond is observed.8 (iv) Metal–nitrogen bound s complexes. In this mode of co-ordination, tcne/q may bind as a neutral or anionic 9 ligand.Species featuring ‘intermediate oxidation states’,10 or containing both the neutral and the anionic tetranitrile,11 have also been reported. The compounds tcne/q may also act as bridging ligands, providing a conjugated link between two or more metal centres.12,13 Ligand geometries tend to be similar to those observed in the analogous oxidation state of the unbound tetranitrile and, in the solid state, ligand ‘stacking’ is often observed.9,10,13,14 (v) Complexes in which tcne has reacted with other ligands.Insertions into metal–ligand bonds (usually into M]H, M]C or M]N bonds), 2 1 2 and 2 1 4 cycloaddition reactions (tcne is an excellent dienophile) and rearrangement reactions are the most commonly reported examples of this phenomenon.15 Occasionally, species are † Present address: Department of Chemistry, University of Nottingham, University Park, Nottingham NG7 2RD, UK. ‡ Supplementary data available (No. SUP 57238, 2 pp.): fast atom bombardment mass spectrometry data for trans-[OsCl2(PEt2Ph)3(tcva)].See Instructions for Authors, J. Chem. Soc., Dalton Trans., 1997, Issue 1. reported in which interactions between the metal centre and tcne/q fall into more than one of the above categories.1,16 Complexes of general formula [OsX3(PR2R9)3] (X = Cl or Br; R, R9 = alkyl or aryl groups and may be equivalent) appear to be well suited to reactions with a strong p-acceptor ligand such as tcne or tcnq.Halide loss is readily induced (usually by reduction of the starting complex), and the resultant species reacts with various compounds usually giving complexes of formula [OsX2(PR2R9)3L]n1 or [OsX2(PR2R9)2L2]n1 (L = dmf, PhCN, MeCN, N2, CO, PR3, etc.).17,18 In general, the product complexes contain OsII. It is possible to induce such reactions by either chemical 17 or electrosynthetic 18 methods, although it has been found that the stereochemistry of the product may depend on the methods employed. In cases where mer- [OsX2(PR2R9)3L]n1 is obtained the trans-X2 isomer is usually the kinetically favoured product, but the cis-X2 isomer tends to have greater thermodynamic stability.18 We report here the reaction of mer-[OsCl3(PEt2Ph)3] with tcne, which results in the formation of the unexpected product trans- [OsCl2(PEt2Ph)3(tcva)] [tcva = tricyanovinyl alcohol (1,1,2-1974 J.Chem. Soc., Dalton Trans., 1997, Pages 1973–1979 tricyano-2-hydroxyethylene)].The structural and electronic characterisation of this product are reported and discussed. Experimental Infrared spectra were recorded on a Perkin-Elmer 598 IR spectrometer, UV/VIS/NIR spectra on a Perkin-Elmer Lambda-9 spectrophotometer, spectroelectrochemical studies being conducted within an optically transparent electrode cell,19 EPR spectra were recorded on a Bruker ER-200D-SRC spectrometer, using an in situ spectroelectrochemical cell.19 In both UV/VIS/NIR and EPR spectroelectrochemical studies, the potential source was a Metrohm E506 potentiostat.Infrared spectroelectrochemical studies involved electrosynthesis in a conventional coulometric cell and transfer of the electrogenerated species to solution IR cells. Voltammetric and coulometric studies were conducted using a PAR-170 electrochemistry system, and were recorded on a Hewlett-Packard 7045A X-Y recorder. All electrochemical experiments were conducted with [NBun 4][BF4] as supporting electrolyte and the solutions were purged with Ar for 30 min prior to study.Electrochemical experiments were performed using a conventional three-electrode configuration, with Pt micro-working and counter electrodes and a Ag–AgCl reference electrode against which the ferrocene–ferrocenium couple was measured at 10.55 V. Cyclic voltammograms were recorded using a scan rate of 100 mV s21, while a.c. and stirred voltammograms were recorded at a scan rate of 20 mV s21. Fast atom bombardment (FAB) mass spectra were recorded on a Kratos-50TC spectrometer, using thioglycerol as the supporting matrix. Compounds OsO4, tcne, PEt2Ph, AgBF4, NBun 4OH, HBF4, silica and HCl were used as supplied, [NBun 4][BF4] was prepared by neutralisation of NBun 4OH with HBF4 in water, the precipitated product being purified by recrystallisation from water and methanol and dried in vacuo.Dichloromethane was purified by standing over KOH pellets, then distillation over P2O5, under N2, tetrahydrofuran (thf) was purified by distillation over Na wire and benzophenone, under N2, hexane was purified by distillation over Na wire, under N2.Other solvents were used as supplied. Preparations mer-[OsCl3(PEt2Ph)3]. This complex was prepared by a method similar to that published by Chatt et al.20 Concentrated HCl (1.0 cm3) and PEt2Ph (2.00 g, 12.0 mmol) were added, under N2, to nitrogen-purged ethanol. The compound OsO4 (1.00 g, 3.93 mmol) was added, and the reaction mixture was heated to reflux for 15 min.After cooling, crude mer- [OsCl3(PEt2Ph)3] was obtained by filtration {the main impurity being [(PEt2Ph)3Os(m-Cl)3Os(PEt2Ph)3]1}. The product was purified by elution with CH2Cl2 from a silica column, in 60% yield. FAB mass spectrum: m/z 795, [C30H45Cl3OsP3]1 requires m/z 795. IR (KBr): 1482m, 1460 (sh), 1451m, 1432s, 1250m, 1100m, 1042s, 1030s, 762m, 752m, 745m, 728s, 704s, 698s cm21. trans-[OsCl2(PEt2Ph)3(tcva)]. The complex mer-[OsCl3(PEt2- Ph)3] (140 mg, 0.176 mmol) and tcne (45 mg, 0.351 mmol) were dissolved in nitrogen-purged thf (30 cm3), under N2.A solution of AgBF4 (40 mg, 0.21 mmol) in thf (10 cm3) was added, inducing immediate reaction. The reaction mixture was allowed to stand in darkness while the AgCl precipitate settled, then filtered under gravity. The volume of solvent was reduced to approximately 10 cm3 and the product was precipitated by addition of hexane (25 cm3), filtered off and washed repeatedly with hexane.The yield of trans-[OsCl2(PEt2Ph)3(tcva)] was 120 mg (78%) (Found: C, 48.2; H, 5.3; N, 4.9. Calc. for C35H46Cl2N3OOsP3: C, 47.8; H, 5.25; N, 4.8%). FAB mass spectrum: m/z 879, 843, 760, 725, 594, 555, 525, 429, 393, 167 (SUP 57238). IR (KBr): 3250m, 2225 (sh), 2220m, 2185m, 1610s, 325m, 310w cm21. Overlaying a thf solution of the complex with n-hexane resulted in formation of crystals suitable for X-ray diffraction. Crystallography Crystallographic data for trans-[OsCl2(PEt2Ph)3(tcva)]?0.5- C4H8O. Deep red lath of dimensions 0.16 × 0.25 × 0.54 mm; C37H50Cl2N3O1.5OsP3; M = 919.9.The space group was determined from q values of 23 reflections measured at ±w such that 14 < q < 168. Monoclinic, space group = P21/n (alternate no. 14 of P21/c), a = 12.8579(14), b = 18.313(2), c = 17.144(2) Å, b = 94.349(11)8, Z = 4, Dc = 1.509 g cm23, m = 34.55 cm21, F(000) = 1840. Collection and processing of crystallographic data.The crystal was fixed to a glass fibre and mounted on a Stoe Stadi-4 four-circle diffractometer. Data were collected at 298 K in the w–2q mode using graphite-monochromated Mo-Ka radiation (l = 0.710 74 Å). Of 5446 unique reflections measured (2.5 < q < 22.5; h = 213 to 13, k = 0–18 and l = 0–18), 4593 data with F > 4s(F) were used in all calculations. An initial correction for absorption was made21 (maximum and minimum transmission factors 0.4207 and 0.3102 respectively).Structure solution and refinement were carried out using SHELX 76,22 the Os atom of the asymmetric unit being located by Patterson synthesis. All other non-hydrogen atoms were located by a process of iterative least-squares refinement (based on F) and Fourier-difference synthesis. Scattering factors for Os were obtained from ref. 23, all other scattering factors were inlaid in SHELX 76. All non-hydrogen atoms except for those within the cocrystallised solvent were refined anisotropically. The C atoms of the phenyl rings were constrained as regular hexagons.The H atom of tricyanovinyl alcohol was located, then refined isotropically with the O]H distance restrained. All other H atoms were refined isotropically in fixed, calculated positions. At isotropic convergence, a final absorption correction was made using DIFABS24 (maximum and minimum corrections 1.188 and 0.913 respectively). The weighting scheme w21 = s2(F) 1 0.000244F2 was found to give satisfactory agreement analyses.At convergence, the final values of R and R9 were 0.026 and 0.038 respectively for 393 parameters, with a goodness of fit index of 1.234. In the final Fourierdifference synthesis the maximum and minimum residual electron densities were 0.58 and 20.44 e Å23 respectively. Results and Discussion Upon addition of AgBF4 to solutions of tcne and mer- [OsCl3(PEt2Ph)3], in thf or CH2Cl2, an immediate colour change from orange to blue-grey is observed.In thf, the reaction mixture then turns red-purple over the next 30 min as trans-[OsCl2(PEt2Ph)3(tcva)] is formed. However, in CH2Cl2 it is possible to interrupt the reaction at the intermediate stage. The blue-grey intermediate proved to be extremely reactive and has not been fully characterised. Structural characterisation Fast atom bombardment mass spectrometry data for trans- [OsCl2(PEt2Ph)3(tcva)] shows the molecular ion and various fragments resulting from its breakdown.The elemental analytical data are in good agreement with the calculated composition of trans-[OsCl2(PEt2Ph)3(tcva)]. Infrared spectra of trans-[OsCl2(PEt2Ph)3(tcva)], recorded as doped KBr discs, show a number of bands in addition to those associated with PEt2Ph and two Os]Cl stretches at 325 and 310 cm21. Three CN stretching modes are observed at 2225 (sh), 2220m and 2185m cm21. More importantly, a broad band corresponding to an O]H stretching vibration is observed at 3250m cm21.This confirms that the ligand is tricyanovinyl alcohol and not its deprotonated anion. A strong bandJ. Chem. Soc., Dalton Trans., 1997, Pages 1973–1979 1975 observed at 1610 cm21 was tentatively assigned as the C]] C stretching mode in tcva [in free tcne the C]] C stretch is observed in Raman spectra at 1569 cm21 (ref. 25)]. It was found that the energies of the CN stretching modes in trans-[OsCl2- (PEt2Ph)3(tcva)] are unaffected by dissolution of the complex in CH2Cl2.A single-crystal X-ray diffraction study of trans- [OsCl2(PEt2Ph)3(tcva)] confirmed its structure, structural data are listed in Table 1. Views of trans-[OsCl2(PEt2Ph)3(tcva)] are shown in Fig. 1. The complex cocrystallises with thf, with an approximate stoichiometry of trans-[OsCl2(PEt2Ph)3(tcva)]?0.5 thf. There are four molecules of trans-[OsCl2(PEt2Ph)3(tcva)] per monoclinic unit cell. As a result of disorder, no distinction has been made between oxygen and the carbon atoms of the cocrystallised thf, and the hydrogen atoms are omitted.This is the first reported example of metal-bound tricyanovinyl alcohol. The hydroxy group is found to be bound to the olefinic carbon furthest from the metal, cis to the Os-bound nitrogen. The structure exhibits a number of other noteworthy features. The Os]N bond is fairly long at 2.064(4) Å, in comparison with other N-bound tcne complexes of osmium. For example, in the complexes [Os(S2PR2)2(PPh3)(tcne)], Os]N = 1.899(7) Å and 1.858(10) Å for R = Me and Ph respectively. 10 The relatively long Os]N bond in trans-[OsCl2- (PEt2Ph)3(tcva)] is consistent with relatively weak p back bonding from Os to tcva. This is thought to be due, in part, to tcva being trans to a phosphine ligand, as opposed to the p donor ligands (such as dithiolates) which occupy the analogous positions in [Os(S2PR2)2(PPh3)(tcne)]. The alkene C]] C bond length of the tcva ligand is 1.328(9) Å, which is not significantly different from that observed in free tcne [1.344(3) Å26].The short C]] C bond length therefore supports the assignment of the band at 1610 cm21 in the IR spectrum to the alkene stretching mode in tcva. The C]] C bond length is also consistent with observations regarding the long Os]N bond length and relatively low p back bonding. It has been shown for tcne complexes that weakening of the C]] C bond is related to the extent of back bonding into the tcne p* lowest unoccupied molecular orbital (LUMO).1,9,10 Thus, the reported C]] C bond in [Fe(C5Me5)2][tcne] is 1.392(8) Å,3 whereas the alkene bond length has increased to 1.49(2) Å in [{Co- (C5Me5)2}2][tcne].2 The weakening of the C]] C bond in such complexes is most effectively monitored using IR spectroscopy, where the shift of the C]] C stretching frequency has been used to assign N-bound tcne complexes as containing anionic tcne.1,9 Crystallographic data for s-bound tcne complexes, however, Table 1 Selected bond lengths (Å) and angles (8) of trans- [OsCl2(PEt2Ph)3(tcva)] (estimated standard deviations in parentheses) Os]N(11) Os]Cl(1) Os]Cl(2) Os]P(1) Os]P(2) Os]P(3) N(11)]C(11T) 2.064(4) 2.3488(11) 2.3448(12) 2.4270(12) 2.3763(12) 2.4501(12) 1.151(6) C(11T)]C(1T) C(1T)]C(12T) C(1T)]C(2T) C(12T)]N(12) C(2T)]O(1) C(2T)]C(21T) C(21T)]N(22) 1.406(7) 1.451(9) 1.328(9) 1.056(12) 1.178(10) 1.534(12) 1.160(12) N(11)]Os]Cl(1) N(11)]Os]P(1) N(11)]Os]P(3) Cl(1)]Os]P(1) Cl(1)]Os]P(3) Cl(2)]Os]P(2) P(1)]Os]P(2) P(2)]Os]P(3) N(11)]C(11T)]C(1T) C(11T)]C(1T)]C(2T) C(1T)]C(12T)]N(12) C(1T)]C(2T)]C(21T) C(2T)]C(21T)]N(22) 87.92(10) 85.32(10) 85.09(10) 88.45(4) 87.54(4) 87.75(4) 94.99(4) 94.75(4) 176.7(5) 123.8(5) 176.4(9) 111.8(6) 176.9(9) N(11)]Os]Cl(2) N(11)]Os]P(2) Cl(1)]Os]Cl(2) Cl(1)]Os]P(2) Cl(2)]Os]P(1) Cl(2)]Os]P(3) P(1)]Os]P(3) Os]N(11)]C(11T) C(11T)]C(1T)]C(12T) C(12T)]C(1T)]Cl(2T) C(1T)]C(2T)]O(1) O(1)]C(2T)]C(21T) 89.87(10) 177.60(10) 177.66(4) 94.46(4) 92.13(4) 91.51(4) 169.73(4) 173.6(4) 114.3(5) 121.6(6) 126.7(7) 121.4(7) tend to be rather less informative with regard to the relationship between p back bonding and C]] C bond lengths.The primary problem is the large uncertainty usually associated with the relevant distance in the reported complexes. For example, in [Os(S2PPh2)2(PPh3)(tcne)], the tcne C]] C bond length was found to be 1.401(19) Å,10 a distance which, because of the large uncertainty associated with it, cannot be said to differ signifi- cantly from the corresponding bond length in free tcne.26 The IR spectrum of [Os(S2PPh2)2(PPh3)(tcne)], however, contains a C]] C stretch at 1430 cm21,10 indicating significant weakening of the olefinic bond.The C]O bond in trans-[OsCl2(PEt2Ph)3(tcva)] is unusually short for a single bond, at 1.178(10) Å. However, this is consistent with structural data for salts of the deprotonated anion of tcva (structural data are not available for free tcva), with reported C]O bond lengths in the range 1.15(2)–1.246(7) Å.27 The short C]O distance suggests significant interaction between oxygen and the olefinic p orbitals.The phenyl groups bound to P(1) and P(2) lie in planes which are parallel to within 1.68. The perpendicular separation of the rings is 3.3 Å and the centre–centre distance is 3.8 Å, suggesting van der Waals interactions similar to those encountered in graphite. Such structural phenomena have been observed previously in arylphosphine complexes of transition metals, e.g.[OsCl2(PMe2Ph)4].28 The P atom trans to tcva [P(2)] is also significantly closer to Os than P(1) and P(3), possibly as a result of the interaction between the phenyl rings. This proximity is further evidence that tcva is acting as a weak p acceptor in trans-[OsCl2(PEt2Ph)3(tcva)]. Ligand stacking, a fairly common phenomenon in tcne and tcnq complexes,1,9,10,13,14 is not observed in the crystal structure of trans-[OsCl2(PEt2Ph)3(tcva)]. Electronic characterisation In trans-[OsCl2(PEt2Ph)3(tcva)] there are two redox-active sites, the metal centre and the co-ordinated tcva.Given that the related tcne may co-ordinate to a metal centre as a neutral or an anionic ligand, it is clearly important that the interactions of these redox sites be characterised. The structural data suggest that weak electron donation from Os to tcva occurs in trans- [OsCl2(PEt2Ph)3(tcva)] [i.e.the complex may be described as Fig. 1 Views of trans-[OsCl2(PEt2Ph)3(tcva)]1976 J. Chem. Soc., Dalton Trans., 1997, Pages 1973–1979 Table 2 Data from spectroelectrochemical studies of trans-[OsCl2(PEt2Ph)3(tcva)] Spectroscopic technique Species IR, n& /cm21 EPR UV/VIS, n/cm21 trans-[OsCl2(PEt2Ph)3(tcva)] trans-[OsCl2(PEt2Ph)3(tcva)]2 trans-[OsCl2(PEt2Ph)3(tcva)]1 2225, 2220, 2185 2180, 2140, 2105 No signal observed g = 2.012, D2� 1 = 23 G No signal observed 17 400, 20 400, 32 500, 39 000 24 000, 34 000, 39 500 13 900, 23 700, 36 000 OsII(tcva)0, as opposed to OsIII(tcva)2], but electrochemical and spectroelectrochemical studies are required for a rigorous characterisation of the frontier orbitals.Redox behaviour. The cyclic voltammogram of tcne reveals two reductions. In CH2Cl2, at room temperature, the first reduction is fully reversible and observed at 10.40 V vs. Ag–AgCl. The second reduction, at 20.7 V, is only quasi-reversible at room temperature.Unfortunately, no redox data for free tricyanovinyl alcohol are available. The cyclic voltammogram of mer-[OsCl3(PEt2Ph)3], in CH2Cl2 at room temperature, reveals a reversible OsIII/IV oxidation at 11.05 V, and a partially reversible OsIII/II reduction at 20.40 V. The chemically irreversible reduction is associated with the appearance of a redox process at 10.22V which corresponds to the oxidation of the five-coordinate intermediate. Fig. 2 shows cyclic and a.c.voltammograms of trans- [OsCl2(PEt2Ph)3(tcva)] in CH2Cl2 at room temperature. A reversible oxidation is observed at 11.43 V, along with a reversible reduction at 10.09 V and an irreversible reduction at 21.0 V. Coulometric studies reveal that the reversible reduction and oxidation couples are both one-electron processes. On replacement of a p donor such as Cl2 with a p acceptor ligand, one would predict a large anodic shift in the potential of any metal-based redox processes.It is therefore logical to assign the oxidation at 11.43 V as metal-based. Although redox data are not available for free tcva we suggest that the redox potentials will be comparable with those of tcne. Therefore, we provisionally assign the two reduction processes to be sequential one-electron reductions to the LUMO of tcva. The voltam- Fig. 2 Cyclic and a.c. voltammograms of trans-[OsCl2(PEt2Ph)3(tcva)] in CH2Cl2 at room temperature metric data are therefore generally consistent with relatively weak p back donation from Os into the tcva LUMO.Spectroelectrochemical studies. As a means of validating the above assignments, the spectroelectrochemical behaviour of trans-[OsCl2(PEt2Ph)3(tcva)] was investigated. Complete spectroelectrochemical data are quoted in Table 2. Infrared spectra of trans-[OsCl2(PEt2Ph)3(tcva)] and its monoanion were recorded in CH2Cl2 and show a significant shift to lower energy of the CN stretching modes of tcva, from 2225, 2220 and 2185 cm21 to 2180, 2140 and 2105 cm21 respectively.The CN stretching vibrations of free tcne and tcne2 in CH2Cl2, were observed at 2260 and 2215, and at 2180 and 2140 cm21 respectively. Clearly, there is a significant weakening in the CN bonds on reduction of trans-[OsCl2(PEt2Ph)3(tcva)]. In addition, EPR spectroelectrochemical studies were also conducted, trans-[OsCl2(PEt2Ph)3(tcva)] and its first oxidation product gave no observable signal down to 77 K.The frozenglass spectrum of trans-[OsCl2(PEt2Ph)3(tcva)]2, in MeCN, gave a single line signal where g = 2.010 and the peak width = 23 G. This result is consistent with the unpaired electron being localised on tcva. Attempts to resolve coupling to the 14N nuclei of tcva were unsuccessful. The lack of an observable EPR response from trans-[OsCl2(PEt2Ph)3(tcva)]1 may be due to a multiply degenerate (or nearly degenerate) electronic ground state, resulting in extremely rapid relaxation when the unpaired electron is excited.This behaviour is not inconsistent with the unpaired electron being in an Os-based orbital. On the basis of the IR and EPR spectroelectrochemical studies, and the voltammetric analysis, it is reasonable to conclude that the complex under study is best described as trans-[OsIICl2(PEt2Ph)3(tcva)0], which may be oxidised to trans-[OsIIICl2(PEt2Ph)3(tcva)0]1, and reduced to trans- [OsIICl2(PEt2Ph)3(tcva)2]. Fig. 3 The UV/VIS absorption spectrum of mer-[OsCl3(PEt2Ph)3]- in CH2Cl2 at 290 KJ.Chem. Soc., Dalton Trans., 1997, Pages 1973–1979 1977 The UV/VIS spectroelectrochemistry of trans-[OsCl2- (PEt2Ph)3(tcva)] was also investigated. The electronic absorption spectrum of mer-[OsCl3(PEt2Ph)3] in CH2Cl2 at room temperature, is shown in Fig. 3. Bands are observed at 18 900 (e 700 dm3 mol21 cm21), 23 100 (1700), 29 500 (sh, 1400), 33 400 (2600) and 39 500 cm21 (13 500). A similar spectrum was reported for mer-[OsCl3(PBu2Ph)3], in which the highest energy band was assigned to a Cl2 æÆ Os(eg*) process and the other four were assigned as Cl2 æÆ Os(t2g) charge-transfer bands.29 The assignment of the highest energy band is, however, called into question by the spectra of two related species.In the spectrum of mer-[OsBr3(PEt2Ph)3] the four lower energy bands have red shifted to 17 300, 19 200, 20 700 and 24 900 cm21, but the highest energy band of the five was unaffected by the change of halide.29 This indicates that the highest energy band does not arise from a halide–metal charge-transfer process.Furthermore, in the spectrum of mer-[OsCl3(PPrn 3)3], the four lower energy bands are observed at similar energies to the above Fig. 4 The UV/VIS spectra showing the reduion of tcne to tcne2 at 233 K in CH2Cl2, Eapp = 10.2 V vs. Ag–AgCl Fig. 5 The UV/VIS absorption spectrum of trans-[OsCl2(PEt2Ph)3- (tcva)] in CH2Cl2 at 260 K trichloro complexes [19 100, 23 900, 26 100 (sh) and 32 700 cm21] but the fifth band is absent.29 We therefore assign the band at 38 500 cm21 to an intraligand transition within the phenyl group of the dialkylaryl phosphine, and not to a ligandmetal charge transfer (l.m.c.t.) process.The electronic absorption spectrum of tcva has not been reported, but it is reasonable to expect that there will be similarities to the spectrum of tcne. Reduction of tcne to the tcne2 radical anion was conducted within the optically transparent electrode cell at 233 K, in CH2Cl2.Upon reduction (shown in Fig. 4) the p æÆ p* band at 39 000 cm21 moves to 23 000 cm21. Vibrational fine structure may be observed in the spectra of both neutral and monoanionic tcne. These observations are consistent with previously published results.30 Note that on completion of the spectroelectrochemical experiment the reduction was reversed and the spectrum of the starting material was fully restored. This confirms that the chemical integrity of the solution remains intact throughout the experiment.The spectrum of trans-[OsCl2(PEt2Ph)3(tcva)] in CH2Cl2 at 260 K, is shown in Fig. 5. Maxima were found at 17 400 (e 600 dm3 mol21 cm21), 20 400 (900), 32 500 (5500) and 39 000 cm21 (14 200). The two lower energy bands were found to exhibit significant solvatochromism: in thf, for example, they are found at 18 100 and 20 900 cm21. Such solvent dependence indicates that the transitions associated with the lower energy bands have a significant charge-transfer component.Conversely, the positions of the two highest energy bands are unaffected by changes of solvent. The band at 39 000 cm21 was again assigned to an intraligand transition on the phosphine ligand. Given the results of the IR and EPR spectroelectrochemical studies, one would predict the presence of OsII æÆ tcva metal-ligand charge-transfer (m.l.c.t.) transitions at fairly low energy, the bands at 17 400 and 20 400 cm21 were therefore assigned to such m.l.c.t.processes. In light of its solvent independence, the band at 32 500 cm21 was also assigned to an intraligand transition, presumably tcva-based. Reduction of trans-[OsCl2(PEt2Ph)3(tcva)] in CH2Cl2 at 260 K, is shown in Fig. 6. Collapse of the lowest energy band was observed, while a new band appeared at 24 000 cm21. Changes in the UV range of the spectrum were relatively minor. Several isosbestic points were observed, indicating a clean electron- Fig. 6 The UV/VIS spectra showing reduction of trans- [OsCl2(PEt2Ph)3(tcva)] to trans-[OsCl2(PEt2Ph)3(tcva)]2 in CH2Cl2 at 260 K, Eapp = 20.1 V1978 J. Chem. Soc., Dalton Trans., 1997, Pages 1973–1979 transfer reaction. We assign the band at 24 000 cm21 to an intraligand transition of tcva2, by analogy with the spectrum of tcne2. Oxidation of trans-[OsCl2(PEt2Ph)3(tcva)] also appears to be a one-step process, readily reversible at 260 K. Isosbestic points are again observed, this oxidation is shown in Fig. 7. The spectrum of trans-[OsCl2(PEt2Ph)3(tcva)]1 is found to contain maxima at 13 900 (e 1300 dm3 mol21 cm21), 23 700 (1200) and 36 000 (14 000). The highest energy band was assigned to an intraligand p æÆ p* transition on the phosphine ligand. The band at 13 900 cm21 was assigned to a tcva æÆ OsIII charge-transfer process, while that at 23 700 cm21 has been assigned as a Cl2 æÆ OsIII l.m.c.t. band. Conclusion To conclude, we have reported the synthesis and characterisation of an osmium complex featuring a novel, redox-active ligand.Structural, electrochemical and spectroelectrochemical investigations indicate that the highest occupied molecular orbital (HOMO) is metal-based, while the LUMO is ligandbased. We formulate the neutral species as trans-[OsIICl2- (PEt2Ph)3(tcva)0], its oxidation product as trans-[OsIIICl2- (PEt2Ph)3(tcva)0]1, and the reduced complex as trans-[OsIICl2- (PEt2Ph)3(tcva)2]2. Preliminary mechanistic studies lead us to believe that the substitution reaction which yields tricyanovinyl alcohol is the result of small amounts of water present in thf.It is, for example, possible to induce formation of trans-[OsCl2- (PEt2Ph)3(tcva)] in CH2Cl2 by the simple expedient of adding a small amount of water to the reaction mixture, either before or after formation of the blue-grey intermediate. Replacement of CN by OH is only observed cis to the Os-bound nitrogen, strongly suggesting that substitution occurs after co-ordination of tcne to osmium.Furthermore, our current data on the bluegrey intermediate [IR (KBr) 2215 (sh), 2205 (sh), 2195m cm21, no OH stretch; UV/VIS (CH2Cl2) 13 800 (1400), 39 000 cm21 (14 000 dm3 mol21 cm21); reversible oxidation at 11.52 V and reversible reduction at 20.19 V vs. Ag–AgCl] suggest that it is trans-[OsCl2(PEt2Ph)3(tcne)]. Unanswered questions remain, however, regarding the mech- Fig. 7 The UV/VIS spectra showing oxidation of trans- [OsCl2(PEt2Ph)3(tcva)] to trans-[OsCl2(PEt2Ph)3(tcva)]1 in CH2Cl2 at 260 K, Eapp = 11.8 V anism by which trans-[OsCl2(PEt2Ph)3(tcva)] is formed.The identity of the blue-grey intermediate requires confirmation and the product formulation indicates that the metal centre has been reduced. In the original preparation, AgBF4 was added as a chloride abstractor. However, experiments using TlBF4 give no reaction, whereas bulk electrochemical reduction of the starting materials in a coulometric cell, in CH2Cl2, gives the unstable blue-grey intermediate noted previously. We therefore conclude that AgBF4 is more intimately involved in the reaction than as a simple chloride abstractor.Mechanistic studies are currently being undertaken. The formation of trans-[OsCl2- (PEt2Ph)3(tcva)] also suggests the possibility of conducting reactions of Os-bound tcne with other nucleophiles. Acknowledgements We thank the EPSRC for financial support.References 1 W. Kaim and M. Moscherosch, Coord. Chem. Rev., 1994, 129, 157. 2 D. A. Dixon and J. S. Miller, J. Am. Chem. Soc., 1987, 109, 3656. 3 J. S. Miller, J. C. Calabrese, S. R. Chittapeddi, J. H. Zhang, W. H. Reiff and A. J. Epstein, J. Am. Chem. Soc., 1987, 109, 769. 4 J. S. Miller, D. M. O’Hare, A. Chakraborty and A. J. Epstein, J. Am. Chem. Soc., 1989, 111, 7853; W. E. Broderick, J. A. Thompson, E. P. Day and B. M. Hoffmann, Science, 1990, 249, 401; K.-M.Chi, J. C. Calabrese, W. M. Reiff and J. S. Miller, Organometallics, 1991, 10, 688; G. A. Gandela, J. S. Miller and M. J. Rice, J. Am. Chem. Soc., 1979, 101, 2755. 5 P. Kathirgamanathan and D. R. Rosseinsky, J. Chem. Soc., Chem. Commun., 1980, 839; H. Kitagawa, T. Mitani, J. Toyoda, K. Nakasuji, H. Okamoto and M. Yamashita, Synth. Met., 1993, 56, 1783; V. J. Murphy and D. O’Hare, Inorg. Chem., 1994, 33, 787; P. Zhou, J. S. Miller and A. J. Epstein, Phys. Lett. A, 1994, 189, 193. 6 M. Rosenblum, R. W. Fish and C. Bennet, J. Am. Chem. Soc., 1964, 86, 5166. 7 M. D. Ward, P. J. Fagan, J. C. Calabrese and D. C. Johnson, J. Am. Chem. Soc., 1989, 111, 1719; M. D. Ward, Synth. Met., 1988, 27, B211. 8 J. A. McGinnetty and J. A. Ibers, J. Chem. Soc., Chem. Commun., 1968, 235; G. Mestroni, A. Camus and G. Zassinovich, J. Organomet. Chem., 1974, 65, 119; O. Gandolfi, B. Giovannitti, M. Ghedini and G. Dolcetti, J. Organomet. Chem., 1976, 129, 207; T. S. Janik, K. A.Bernard, M. R. Churchill and J. D. Atwood, J. Organomet. Chem., 1987, 323, 247; M. Bottrill, M. Green, A. G. Orpen, D. R. Saunders and I. D. Williams, J. Chem. Soc., Dalton Trans., 1989, 511; J. K. Stalick and J. A. Ibers, J. Am. Chem. Soc., 1970, 92, 5333; A. Maissonat, J.-J. Bonnet and R. Poilblanc, Inorg. Chem., 1980, 19, 3168. 9 M. F. Rettig and R. M. Wing, Inorg. Chem., 1969, 8, 2685; R. Gross and W. Kaim, Angew. Chem., Int. Ed. Engl., 1987, 26, 251. 10 A. E. D. McQueen, A.J. Blake, T. A. Stephenson, M. Schröder and L. J. Yellowlees, J. Chem. Soc., Chem. Commun., 1988, 1533; A. E. D. McQueen, Ph.D. Thesis, University of Edinburgh, 1988. 11 H. Oshio, E. Ino, T. Ito and Y. Maeda, Bull. Chem. Soc. Jpn., 1995, 68, 889. 12 S. L. Bartley and K. R. Dunbar, Angew. Chem., Int. Ed. Engl., 1991, 30, 448. 13 R. Gross-Lannert, W. Kaim and B. Olbrich-Deussner, Inorg. Chem., 1990, 29, 5046; W. Kaim, T. Roth, B. Olbrich-Deussner, R. Gross- Lannert, J. Jordanov and E.K. H. Roth, J. Am. Chem. Soc., 1992, 114, 5693; L. Ballester, M. C. Barral, A. Gutierrez, R. Jiminez- Aparicio, J. M. Martinez-Muyo, M. F. Perpinan, M. A. Monge and C. Ruiz-Valero, J. Chem. Soc., Chem. Commun., 1991, 1396; L. Ballester, M. C. Barral, R. Jimenez-Aparicio and B. Olombrada, Polyhedron, 1996, 15, 218. 14 H. Braunwarth, G. Huttner and L. Zsolnai, J. Organomet. Chem., 1989, 372, C23; J. P. Cornelissen, J. H. van Diemen, L. R. Groenveld, J. G. Hasnoot, A.L. Spek and J. Reedijk, Inorg. Chem., 1992, 31, 199; B. Olbrich-Deussner, W. Kaim and R. Gross-Lannert, Inorg. Chem., 1989, 28, 3113. 15 M. I. Bruce, T. W. Hambley, M. R. Snow and A. G. Swincer, Organometallics, 1985, 4, 501; K. L. Amos and N. G. Connelly, J. Organomet. Chem., 1980, 194, C57. 16 M. M. Olmstead, G. Speier and L. Szabo, J. Chem. Soc., Chem. Commun., 1994, 541.J. Chem. Soc., Dalton Trans., 1997, Pages 1973–1979 1979 17 J. Chatt, G. J. Leigh and R. L. Richards, J. Chem. Soc. A, 1970, 2243; J. Chatt, D. P. Melville and R. L. Richards, J. Chem. Soc. A, 1971, 1169; J. Chatt, G. J. Leigh and R. M. Paske, J. Chem. Soc. A, 1969, 854; B. Bell, J. Chatt, J. R. Dilworth and G. J. Leigh, Inorg. Chim. Acta, 1972, 6, 635; D. J. Cole-Hamilton and T. A. Stephenson, J. Chem. Soc., Dalton Trans., 1976, 2396; G. J. Leigh, J. J. Levison and S. D. Robinson, Chem. Commun., 1969, 705. 18 V. T. Coombe, G. A. Heath, T. A. Stephenson, G. A. Whitelock and L. J. Yellowlees, J. Chem. Soc., Dalton Trans., 1985, 947; V. T. Coombe, Ph.D. Thesis, University of Edinburgh, 1985. 19 S. A. Macgregor, E. McInnes, R. J. Sorbie and L. J. Yellowlees, Molecular Electrochemistry of Inorganic, Bioinorganic and Organometallic Compounds, eds. A. J. L. Pombeiro and J. A. McCleverty, Kluwer Academic Publishers, Dordrecht, 1993, p. 503. 20 J. Chatt. G. J. Leigh, D. M. P. Mingos and R. M. Paske, J. Chem. Soc. A, 1968, 2636. 21 A. C. T. North, D. C. Phillips and D. S. Mathews, Acta Crystallogr., Sect. A, 1968, 24, 351. 22 G. M. Sheldrick, SHELX 76, a program for crystal structure determination and refinement, University of Cambridge, 1976. 23 D. T. Cromer and J. L. Mann, Acta Crystallogr., Sect. A, 1968, 24, 321. 24 N. Walker and D. Stuart, Acta Crystallogr., Sect. A, 1983, 39, 158. 25 J. J. Hinkel and J. P. Devlin, J. Phys. Chem., 1973, 58, 4750. 26 R. G. Little, D. Pautler and P. Coppens, Acta Crystallogr., Sect. B, 1969, 27, 1493. 27 B. W. Sullivan and B. M. Foxman, Organometallics, 1983, 2, 187; T. Dahl, Acta Chem. Scand., Ser. A, 1983, 37, 353. 28 R. J. Sorbie, Ph.D. Thesis, University of Edinburgh, 1989. 29 G. J. Leigh and D. M. P. Mingos, J. Chem. Soc. A, 1970, 587. 30 D. L. Jeanmaire, M. R. Suchanski and R. P. Van Duyne, J. Am. Chem. Soc., 1975, 97, 1699. Received 11th November 1996; Paper 6/07655F

ISSN:1477-9226    

DOI:10.1039/a607655f

出版商:RSC    

年代:1997

数据来源: RSC

摘要:

DALTON J. Chem. Soc., Dalton Trans., 1997, Pages 1981–1988 1981 Mono-, di-, and tri-nuclear complexes of iron(II) with N,N,N9,N9-tetramethylethylenediamine† Sian C. Davies,a David L. Hughes,a G. JeVery Leigh,*,b J. Roger Sanders a and Jaisa S. de Souza c a Nitrogen Fixation Laboratory, John Innes Centre, Norwich Research Park, Colney, Norwich NR4 7UH, UK b School of Chemistry, Physics and Environmental Science, University of Sussex, Brighton BN1 9QJ, UK c Departamento de Química, Universidade Federal do Paraná, Centro Politécnico, 81531-990 Curitiba-PR, Brasil Iron(II) chloride reacted with N,N,N9,N9-tetramethylethylenediamine (tmen) in tetrahydrofuran (thf) to give the mononuclear complex trans-[FeCl2(tmen)2] 1 and the dinuclear complex [{FeCl(tmen)2}2(m-Cl)2] 2.Both adducts were characterised by microanalyses, magnetic measurements, Mössbauer spectroscopy, and crystal structure analysis. In solution in thf the two species are in an equilibrium affected by temperature and by concentration of tmen. Reaction of either 1 or 2 with Na[BPh4] produced the trinuclear species [Fe3(m-Cl)3(m3-Cl)2(tmen)3][BPh4] 3.This compound is paramagnetic both in the solid state and in solution. These iron complexes are very similar in structure to the corresponding vanadium complexes. We have shown how the bulky diamine N,N,N9,N9-tetramethylethylenediamine (tmen) can assist the formation of trinuclear species of vanadium(II) from mononuclear species such as [VCl2- (tmen)2].1 The mechanism of such transformations has not been established, but we reasoned that it should involve an undetected dinuclear species, such as is well established in vanadium( II) chemistry, though not in the form of a tmen derivative. 2 If this is the case, then it might be possible to form a heterometallic trinuclear species by reaction of suitable monoand di-nuclear species of different metals.The heterometallic species of greatest interest to us would contain both iron(II) and vanadium(II), because both these metals are involved in the recently established vanadium–iron nitrogenase.3 However, iron(II) chloride complexes of tmen are not well known, and the only reference we could find is to an undefined material FeCl2(tmen)n, produced by direct reaction of iron(II) chloride with 5 molar equivalents of tmen.4 It was used as a synthon for an iron(II) hydrocarbyl, [Fe(CH2Ph)2(tmen)], believed to have a distorted tetrahedral structure.We therefore commenced our studies by investigating the reaction of iron(II) chloride with tmen rather more closely. Our search for mixed iron(II)– vanadium(II) species is described elsewhere.5 Results and Discussion Structures Tetramethylethylenediamine reacts with anhydrous iron(II) chloride in 3 : 1 molar ratio in tetrahydrofuran (thf) to produce two complexes, as described in the Experimental section. The colourless compound 1, proved to be [FeCl2(tmen)2], on the basis of crystal structural and micro-analysis.It is highly airand moisture-sensitive, and is stable in solution only in an excess of tmen. Otherwise it forms the very pale green compound 2, which was shown to be a dinuclear species [{FeCl2- (tmen)}2(m-Cl)2], again on the basis of crystal structural and micro-analysis. Complex 2 is usually the first product that crystallises from a warm reaction mixture containing an excess of diamine, while 1 is formed slowly at or below room temperature with incorporation of additional tmen.The formation of 1 occurs at the † Non-SI unit employed: mB ª 9.27 × 10224 J T21. expense of 2, and the equilibrium between them depends upon the temperature and the concentration of base. We have observed 1 a similar equilibrium involving mononuclear and trinuclear species in vanadium(II)–tmen chemistry. Reaction of solutions of either 1 or 2 in thf with Na[BPh4] in a 3 : 1 ratio of iron to tetraphenylborate produces crystals of [Fe3Cl5(tmen)3][BPh4] 3 in high yield.This behaviour is very similar to that observed in the vanadium system, and suggests that the M3Cl5 motif is common, at least among first transition series elements. We have since observed such a structure also in nickel chemistry,6 and a cobalt analogue has just been reported.7 We experienced some problems in characterising the very air-sensitive 3. We were never able to observe a molecular ion in the FAB mass spectrum, in part because of sensitivity to the matrix, but we did observe an apparently characteristic peak at m/z at 553.This ion may well be produced in part by interaction with the matrix and we could not assign a structure to it. It was, however, useful for detecting 3 in mixtures.5 These three iron complexes have some intriguing structural features. They are, of course, not completely without precedent. Compound 1 would appear to be a typical high-spin octahedral iron(II) compound.Compound 2 has a dinuclear precedent in complexes such as [Fe2Cl3(thf)6][SnCl5(thf)], though the bridging is completely different.8 We know of no iron precedent for compound 3, but a tetranuclear species [Fe4Cl8(thf)6] with bridging chlorides is well established.9 Clearly there are classes of structure involving iron, chloride, and a neutral ligand of which examples of the first four members have been established. We now discuss our three new compounds in more detail. Compound 1 crystallises in discrete octahedral molecules lying at centres of symmetry.Although there is some structural disorder in the tmen methylene bridges this was successfully modelled by assuming two possible orientations for the methylene carbon C(2), as shown in Fig. 1, and assigning occupancy factors of 50% to the atoms in each of these positions. This kind of phenomenon is not unusual in complexes with ligands such as tmen containing di(methylene) bridges.1,10 The structural data are given in Table 1.The Fe]Cl bond length, 2.397(1) Å, is unremarkable, being in the range for octahedral high-spin FeII complexes.11 It is significantly shorter, though, than the bonds in FeCl2 (2.53 Å, octahedral microsymmetry),12 and this indicates a substantial increase in the covalency of the bond.1982 J. Chem. Soc., Dalton Trans., 1997, Pages 1981–1988 Because the chloride ions occupy trans positions, the Fe]Cl bonds are not subjected to steric pressure, and, as expected, they are ca. 0.1 Å shorter than in the isostructural vanadium(II) complex, trans-[VCl2(tmen)2], the data for which are also included in Table 1.10 The Fe]Cl(3) vector intercepts the equatorial plane at precisely 908, as indicated by the angles N(1)]Fe]Cl(3) [89.9(1)] and N(2)]Fe]Cl(3) [90.0(1)8]. On the other hand, the N(1)]Fe]N(2) angle [79.2(1)8] deviates considerably from 908, as in the vanadium case, and this is again probably a consequence of the small size of the chelate ring.In 1 the Fe]N(1) and Fe]N(2) distances are identical within the experimental error [2.378(3) and 2.376(3) Å, respectively], and longer than the average V]N bond length in the vanadium homologue [2.319(2) Å].10 This is presumably the result of strain due to co-ordination of the two tmen ligands. The high-spin FeII ion is smaller than VII (ionic radii 0.78 and 0.88 Å, respectively). 13 Even so the metal–nitrogen bonds are longer than in [VCl2(tmen)2], probably in order to accommodate the amine molecules around the metal in the presence of the closer chlorides. The analytical data obtained for complex 2 are consistent with the formulation [Fe3Cl5(tmen)3]Cl, but the conductivity data are not.The crystal structure determination showed finally that ‘FeCl2(tmen)’ is actually the dinuclear [{FeCl(tmen)}2(m-Cl)2]. There are two crystallographically independent (but chemically identical) molecules in the unit cell. One lies on a crystallographic centre of symmetry and is affected by disorder, whereas the other is pseudo-centrosymmetric and free of disorder.Fig. 2 is a representation of the pseudo-centrosymmetric dimer, and Table 2 presents selected molecular dimensions. The ratio of non-centrosymmetric to centrosymmetric molecules in the unit cell is 2 : 1. The disorder in the C(14)]N(13)]C(131)]C(132) portion of the tmen ligand in the centrosymmetric dimer was modelled assigning an occupancy factor of 0.5 to each of the possible orientations of C(14), C(131) and C(132).The tmen bridges assumed the usual conformation and dimensions, as found in the non-centrosymmetric molecule and in 1. Fig. 1 Representation of the molecular structure of trans-[FeCl2- (tmen)2] 1, with the atom numbering scheme. The alternative sites which can be occupied by C(2a) and C(2b) in the disordered tmen bridges are depicted Table 1 Selected molecular dimensions (bond lengths in Å, angles in 8) in trans-[MCl2(tmen)2] (M = Fe 1 or V9), with estimated standard deviations (e.s.d.s) in parentheses* M = Fe Fe]N(1) Fe]N(2) Fe]Cl(3) N(1)]Fe]N(2) N(1)]Fe]Cl(3) N(2)]Fe]Cl(3) 2.378(3) 2.376(3) 2.397(1) 79.2(1) 89.9(1) 90.0(1) M = V V]N(1) V]N(2) V]Cl(1) N(1)]V]N(2) N(1)]V]Cl(1) N(2)]V]Cl(1) 2.318(2) 2.320(2) 2.487(1) 81.44(7) 89.82(5) 90.10(5) * Torsion angles in the disordered tmen ligand in the Fe complex: N(1)]C(1)]C(2a)]N(2) 47.8(9), N(1)]C(1)]C(2b)]N(2) 244.5(11).Each FeII ion in 2 is five-co-ordinate and has distorted trigonal-bipyrimdal geometry, Fig. 2. For example, the coordination sphere around the Fe(3) is formed by a bridging chloride, Cl(31), and a nitrogen atom, N(33), in the axial positions, and by the remaining tmen nitrogen, N(36), a second m-Cl, Cl(21), and the terminal chloride, Cl(32), in the equatorial plane. There is considerable distortion in both the axial and equatorial angles, due to the steric strain in the five-membered chelate rings.The Fe ? ? ?Fe distances in the crystallographically distinct molecules are 3.726(2) and 3.732(1) Å, very long for five-coordinate diiron(II) complexes containing bridging chloride ions.14 We found a comparable separation only for [{NiFe- (L)Cl2}2] (3.725 Å) where H2L = N,N9-bis(sulfanylethyl)-1,5- diazacyclooctane 15 in which the square-pyramidal high-spin iron(II) centres are antiferromagnetically coupled. The Fe ? ? ?Fe Fig. 2 Molecular structure of the non-centrosymmetric molecule of [{FeCl(tmen)}2(m-Cl)2] 2, showing the atom numbering scheme Table 2 Selected molecular dimensions (bond lengths in Å, angles in 8) in [{FeCl(tmen)}2(m-Cl)2] 2 with e.s.d.s in parentheses * Molecule 1, centrosymmetric Fe(1) ? ? ?Fe(19) Fe(1)]Cl(119) Fe(1)]Cl(11) Cl(11)]Fe(1)]Cl(12) Cl(11)]Fe(1)]N(13) Cl(11)]Fe(1)]N(16) Cl(12)]Fe(1)]N(13) Cl(12)]Fe(1)]N(16) Fe(1)]Cl(11)]Fe(19) 3.726(2) 2.345(2) 2.648(2) 91.0(1) 170.1(2) 90.7(1) 97.5(1) 120.2(1) 96.4(1) Fe(1)]Cl(12) Fe(1)]N(13) Fe(1)]N(16) N(13)]Fe(1)]N(16) Cl(11)]Fe(1)]Cl(119) Cl(119)]Fe(1)]Cl(12) Cl(119)]Fe(1)]N(13) Cl(119)]Fe(1)]N(16) 2.266(2) 2.286(5) 2.161(5) 80.8(2) 83.7(1) 127.3(1) 95.0(2) 112.3(2) Molecule 2, non-centrosymmetric Fe(2) ? ? ?Fe(3) Fe(2)]Cl(21) Fe(2)]Cl(31) Fe(2)]Cl(22) Fe(2)]N(23) Fe(2)]N(26) Cl(21)]Fe(2)]Cl(31) Cl(22)]Fe(2)]Cl(21) Cl(22)]Fe(2)]Cl(31) Cl(22)]Fe(2)]N(23) Cl(22)]Fe(2)]N(26) N(23)]Fe(2)]N(26) N(23)]Fe(2)]Cl(21) N(23)]Fe(2)]Cl(31) N(26)]Fe(2)]Cl(21) N(26)]Fe(2)]Cl(31) Fe(2)]Cl(21)]Fe(3) 3.732(1) 2.650(2) 2.341(2) 2.264(2) 2.280(5) 2.172(5) 84.0(1) 91.0(1) 129.1(1) 97.0(1) 118.7(2) 81.4(2) 169.9(2) 95.7(1) 89.4(1) 111.9(2) 96.3(1) Fe(3)]Cl(31) Fe(3)]Cl(21) Fe(3)]Cl(32) Fe(3)]N(33) Fe(3)]N(36) Cl(31)]Fe(3)]Cl(21) Cl(31)]Fe(3)]Cl(32) Cl(32)]Fe(3)]Cl(21) Cl(32)]Fe(3)]N(33) Cl(32)]Fe(3)]N(36) N(33)]Fe(3)]N(36) N(33)]Fe(3)]Cl(21) N(33)]Fe(3)]Cl(31) N(36)]Fe(3)]Cl(21) N(36)]Fe(3)]Cl(31) Fe(2)]Cl(31)]Fe(3) 2.667(2) 2.354(2) 2.256(2) 2.272(5) 2.158(5) 83.3(1) 91.1(1) 129.6(1) 97.5(1) 122.7(1) 81.2(2) 94.5(1) 170.4(1) 107.4(1) 90.5(1) 96.1(1) * Torsion angles in the tmen ligands: N(13)]C(14a)]C(15)]N(16) 51.0(15), N(13)]C(14b)]C(15)]N(16) 242.4(16), N(23)]C(24)]C(25)] N(26) 52.2(11), N(33)]C(34)]C(35)]N(36) 48.5(12).J. Chem.Soc., Dalton Trans., 1997, Pages 1981–1988 1983 distance in 2 is much larger than the corresponding distance in [Fe2(m-Cl)3(thf)6]1 [3.086(2) Å], which is dinuclear both in solid state and in solution,8 but the latter is cationic with three bridging chlorides.In the well known tetranuclear [Fe4Cl8(thf)6],9 the mean distance between adjacent iron atoms (3.751 Å) is slightly longer than the Fe ? ? ?Fe separation in 2. These separations are, of course, far greater than any Fe ? ? ?Fe bonding distance,16 and it is therefore not surprising that the tetranuclear compound is said to exist only in the solid state, breaking into dinuclear species in thf solution.9b In the cation [Fe2(m-Cl)3(thf)6]1, the six Fe](m-Cl) bond lengths are approximately equal,8 and average at 2.488(9) Å.In contrast, the {Fe2(m-Cl)2} core of 2 is highly asymmetric, with the Fe](m-Cl)ax mean distance (2.655 Å) being much longer than the Fe-(m-Cl)eq (2.347 Å), and both longer than the average Fe]Cl- (terminal) bond length (2.262 Å), as expected. The asymmetry arises because the bridging chloride atoms are axial to one metal atom and equatorial to the other.14,17 Because of the unsymmetrical nature of the bridging and the large metal– metal separations, this kind of molecule has been stated to be comprised of ‘loosely associated monomers’ in the solid state.17 We have magnetochemical evidence (see below) that dissociation does indeed occur in solution.The diamine is not a particularly good electron donor or pacceptor, and is unable to stabilise the four-co-ordinate ‘FeCl2(tmen)’ species. It is also probably less sterically demanding than dippe [1,2-bis(diisopropylphosphino)ethane].The complex [FeCl2(dippe)] 17 is one of the few examples of unequivocally characterised tetrahedral complexes, other than the tetrahalido complexes, of iron(II). A fine balance between electronic and steric properties is probably responsible for the structure of 2, and of the related mononuclear and trinuclear species 1 and 3. In summary, tmen is not bulky enough to stabilise the electron-deficient tetrahedral [FeCl2(tmen)], too bulky to force a strong electronic interaction between metal ions in [FeCl2- (tmen)2], and just appropriate to form dimeric [Fe2Cl4(tmen)2] in the solid state.This molecule is probably also stable in solution in the absence of an excess of tmen or of another base/ nucleophile at T < 30 8C. These features can be exploited in synthetic work. The X-ray crystallographic analysis of 3 revealed four pairs of ions in the monoclinic unit cell, which is composed of [Fe3Cl5(tmen)3]1 cations and [BPh4]2 anions.The shortest intermolecular contact between the cation and the anion is 3.51(1) Å, from C(52) of a tmen ligand to the phenyl carbon atom C(744), Fig. 3 and supplementary data. Distances Fig. 3 View of the [Fe3(m-Cl)3(m3-Cl)2(tmen)3]1 cation in the tetraphenylborate salt 3, with the atom numbering scheme between anions are larger, the shortest being 3.666(8) Å between phenyl carbon atoms. Small intermolecular distances involving hydrogen atoms were also found.Although these values are smaller than the sum of the van der Waals radii for the atoms involved (3.7–4.0 for C? ? ? C and 2.4 Å for H? ? ?H contacts), they do not imply strong interactions, but rather considerable packing ‘pressures’ in the solid state. The trinuclear cation shows the familiar triangulo arrangement seen in the vanadium(II) systems,1 with the planar {Fe3Cl3} core forming the equatorial plane of the molecule and the triply-bridging chlorides occupying the capping positions (Fig. 3). The co-ordination polyhedron around each FeII is a distorted octahedron, the donor atoms being two of the m-Cl, both m3-Cl and the two N atoms of one tmen. A complete view of the cation with the numbering scheme is shown in Fig. 3. Table 3 contains selected interatomic distances and bond angles. Complex 3 is isostructural with its vanadium homologue (Table 4). There are no large differences in the selected dimensions, that differ (at the most) by ca. 0.1 Å.The main distinction between the two complexes appears in the M? ? ?M non-bonding distances and in the M]Clcap bonds, both being longer in the {Fe3Cl5}1 core. As a consequence, and because the M]Cleq bond lengths are very similar in the two complexes, the Table 3 Selected molecular dimensions (bond lengths in Å, angles in 8) in [Fe3Cl5(tmen)3][BPh4] 3 with e.s.d.s in parentheses a Fe(1) ? ? ?Fe(2) Fe(2) ? ? ?Fe(3) Fe(1)]Cl(1) Fe(1)]Cl(2) Fe(1)]Cl(4) Fe(1)]Cl(5) Fe(1)]N(1) Fe(1)]N(2) Fe(3)]Cl(1) Fe(3)]Cl(2) Fe(3)]Cl(3) Cl(1)]Fe(1)]Cl(2) Cl(1)]Fe(1)]Cl(4) Cl(1)]Fe(1)]Cl(5) Cl(1)]Fe(1)]N(1) Cl(1)]Fe(1)]N(2) Cl(2)]Fe(1)]Cl(4) Cl(2)]Fe(1)]Cl(5) Cl(2)]Fe(1)]N(1) Cl(2)]Fe(1)]N(2) Cl(4)]Fe(1)]Cl(5) Cl(4)]Fe(1)]N(1) Cl(4)]Fe(1)]N(2) Cl(5)]Fe(1)]N(1) Cl(5)]Fe(1)]N(2) N(1)]Fe(1)]N(2) Cl(1)]Fe(3)]Cl(2) Cl(1)]Fe(3)]Cl(3) Cl(1)]Fe(3)]Cl(4) Cl(1)]Fe(3)]N(5) Cl(1)]Fe(3)]N(6) Cl(2)]Fe(3)]Cl(3) Cl(2)]Fe(3)]Cl(4) Cl(2)]Fe(3)]N(5) Fe(1)]Cl(1)]Fe(2) Fe(1)]Cl(1)]Fe(3) Fe(2)]Cl(1)]Fe(3) Fe(1)]Cl(2)]Fe(3) Fe(2)]Cl(3)]Fe(3) 3.222(1) 3.236(1) 2.595(1) 2.501(2) 2.513(1) 2.510(2) 2.209(5) 2.179(4) 2.570(1) 2.495(2) 2.453(2) 81.4 b 87.0 b 81.7(1) 179.4(1) 95.4(1) 82.7 b 158.6(1) 98.8(1) 98.4(1) 83.5(1) 93.6(1) 177.4(1) 98.3(1) 96.0(1) 84.0(2) 82.0 b 82.8(1) 86.3 b 178.3(1) 95.2(1) 159.0(1) 81.6 b 98.3(1) 76.8 b 77.9 b 77.7 b 81.1(1) 81.5(1) Fe(1) ? ? ?Fe(3) Fe(2)]Cl(1) Fe(2)]Cl(3) Fe(2)]Cl(4) Fe(2)]Cl(5) Fe(2)]N(3) Fe(2)]N(4) Fe(3)]Cl(4) Fe(3)]N(5) Fe(3)]N(6) Cl(1)]Fe(2)]Cl(3) Cl(1)]Fe(2)]Cl(4) Cl(1)]Fe(2)]Cl(5) Cl(1)]Fe(2)]N(3) Cl(1)]Fe(2)]N(4) Cl(3)]Fe(2)]Cl(4) Cl(3)]Fe(2)]Cl(5) Cl(3)]Fe(2)]N(3) Cl(3)]Fe(2)]N(4) Cl(4)]Fe(2)]Cl(5) Cl(4)]Fe(2)]N(3) Cl(4)]Fe(2)]N(4) Cl(5)]Fe(2)]N(3) Cl(5)]Fe(2)]N(4) N(3)]Fe(2)]N(4) Cl(2)]Fe(3)]N(6) Cl(3)]Fe(3)]Cl(4) Cl(3)]Fe(3)]N(5) Cl(3)]Fe(3)]N(6) Cl(4)]Fe(3)]N(5) Cl(4)]Fe(3)]N(6) N(5)]Fe(3)]N(6) Fe(1)]Cl(4)]Fe(2) Fe(1)]Cl(4)]Fe(3) Fe(2)]Cl(4)]Fe(3) Fe(1)]Cl(5)]Fe(2) 3.248(1) 2.591(2) 2.501(2) 2.552(2) 2.481(2) 2.202(5) 2.180(5) 2.574(1) 2.196(5) 2.197(5) 81.5(1) 86.3(1) 82.3(1) 94.7(1) 177.4(1) 82.5(1) 159.1(1) 97.8(1) 96.1(1) 83.3(1) 179.0(1) 94.3(1) 96.6(1) 100.3(1) 84.7(2) 97.6(2) 83.0(1) 97.3(2) 98.1(2) 95.4(1) 178.2(1) 83.1(2) 79.0 b 79.3 b 78.3 b 80.4 b a Torsion angles in the tmen ligands: N(1)]C(1)]C(2)]N(2) 59.8(7), N(3)]C(3)]C(4)]N(4) 260.4(8), N(5)]C(5)]C(6)]N(6) 36.1(17).b E.s.d. is less than 0.058.1984 J.Chem. Soc., Dalton Trans., 1997, Pages 1981–1988 octahedral geometry about each FeII ion is even more distorted than in the tris-vanadium cation, with smaller Cleq]M]Cleq angles [158.9(2) vs. 162.1(2)8]. The bond lengths and angles involving the N-donor atoms in the tmen ligand are not signifi- cantly affected, unlike in [MCl2(tmen)2], where the Fe]N bond is longer than the V]N bond. In terms of reactivity, the trisiron( II) complex is probably subjected to higher ring strain caused by the longer Fe ? ? ?Fe non-bonding distances, which could result in more facile substitution of the capping ligands or rupture of the trinuclear core in the presence of donors better than Cl2 or tmen.Magnetic and spectral properties Compound 1 is a normal paramagnetic solid with a temperature-independent magnetic moment in the solid state of 5.3 mB. A slight, reversible temperature dependence was observed in solution measurements, which can be explained by a small degree of association, perhaps involving loss of tmen from the co-ordination sphere.The 1H NMR spectrum gave no indication of unexpected broadening, which could have been caused by the presence of ferromagnetic species due to oxidation or decomposition. The magnetic behaviour of 2 is considerably more complicated. The solid-state data are shown in Fig. 4(a). The slight increase in meff as the temperature is lowered from room temperature to 96 K is not really significant and it was found to be independent of field.Thus this behaviour in unexceptional. However, the situation in solution is quite different. When the magnetic susceptibility was first determined in [2H8]thf solution in the range 2100 to 40 8C (173–313 K), the measurements above room temperature produced data points clearly displaced from the straight line drawn through the remaining data. This was reproducible with different samples prepared by an alternative method. A more detailed investigation was then carried out in the range 250 to 150 8C (223– 323 K), with determination at 38 intervals from 25 to 50 8C [Fig. 4(b)]. The plot indicates ferromagnetism because the experimental values of meff (per FeII ion) increase with decreasing temperature. Fig. 4(b) also shows a small, reproducible, discontinuous transition above room temperature. Preliminary data suggest some thermal hysteresis associated with this change,18 which we are investigating further. These results are not particularly easy to explain.Complex 1 is unlikely to be involved, because it is favoured compared to 2 at lower temperatures and there is no extra tmen present. The observed effect might arise from the formation of a mononuclear, five-co-ordinate species such as [FeCl2(tmen)(thf)], stable on the time-scale of the experiment. However, the FeII system still appears to be ferromagnetically coupled at temperatures higher than the transition temperature and although the overall ferromagnetic effect is weak over the whole temperature range, Table 4 Comparison of mean principal bond lengths (Å) and angles (8) in the triangulo-[M3Cl5(tmen)3][BPh4] complexes (M = V1 or Fe) Complex M? ? ?M M]Cleq M]Clcap M]N Clcap ? ? ? Clcap Cleq]M]Cleq Clcap]M]Clcap N]M]N Cleq]M]Clcap M]Cleq]M M]Clcap]M M = V1 3.142(7) 2.500(4) 2.519(6) 2.214(2) 3.497(1) 162.1(2) 87.9(6) 82.8(3) 83.6(2) 77.8(1) 77.1(2) M = Fe 3.235(8) 2.490(8) 2.566(12) 2.194(5) 3.518(2) 158.9(2) 86.5(2) 83.9(5) 82.4(2) 81.0(3) 78.2(4) it can be rationalised only on the basis of some degree of aggregation.The conversion of 2 into a trinuclear complex such as 3 is conceivable. This happens in vanadium(II) systems, but there is no support for this here, neither from the 1H NMR spectra of 2 or 3 at temperatures corresponding to these measurements nor from the temperature dependence of the magnetic moment/ magnetic susceptibility of 3. In the solid state the effective magnetic moment of [Fe3Cl5- (tmen)3][BPh4] 3 increases from 5.4 at room temperature to 5.8 mB at 90 K, indicating some ferromagnetic coupling between the metal centres. The magnetic behaviour in solution was investigated in a three-step experiment (Fig. 5). The measurements were started at 293 K and repeated at 10–20 K intervals down to 173 K. The system was then warmed to 323 K, measurements being made from 278 to 323 K. Finally, the system was cooled again to 273 K. The effective magnetic moment per iron(II) ion increases with decreasing temperature, consistent with ferromagnetic coupling between iron(II) centres.The effect is reproducible though not dramatic, and is similar to the properties recently reported for some bis(m-halogeno) dinuclear FeII complexes.14 No thermal hysteresis was observed in solution in the temperature range 273 < T < 323 K. Field hysteresis, a phenomenon generally associated with the magnetic behaviour of ferromagnetic systems, could not be detected by either of the methods employed, because all measurements were carried out above the Curie temperature for the complex.In the absence of direct Fe]Fe bonds (as shown by the long M? ? ? M distances determined by X-ray crystallography), the ferromagnetic coupling is probably mediated by the orbitals of the bridging halides. For co-ordination compounds with bridged metal atoms such as 2 and 3, ferromagnetism arises when the singly-occupied metal orbitals involved in the magnetic interaction are orthogonal to each other.The doubly-occupied orbitals on the Fig. 4 Plot of the inverse of the atomic magnetic susceptibility (cA 21, m) and effective magnetic moment (meff, s) versus temperature for [{FeCl(tmen)}2(m-Cl)2] 2 (a) in the solid state and (b) in [2H8]thf solution. Diamagnetic correction: 1.342 × 1024 emu/atom FeIIJ. Chem. Soc., Dalton Trans., 1997, Pages 1981–1988 1985 Table 5 Mössbauer parameters for complexes 1, 2 and 3 (recorded at 77 K, referenced against iron foil at 298 K) Compound 1 [FeCl2(tmen)2] 2 [{FeCl(tmen)}2(m-Cl)2] 3 [Fe3Cl5(tmen)3][BPh4] i.s.a/mm s21 1.17(1) 1.17(1) 1.06(1) 1.06(1) 1.14(1) q.s.a/mm s21 3.10(1) 2.78(1) 3.12(1) 2.83(1) 2.00(1) Ga,b/mm s21 0.15(1) 0.15(1) 0.15(1) 0.16(1) 0.22(1) % total area a,c 56(2) 44(2) 39(3) 61(1) 100 a Numbers in parentheses are the errors in the last significant figure. b Half-width at half maxima.c No constraints applied. bridging ligands allows the establishment of orthogonal threeelectron two-centre bonds with parallel spins on the metal ions.19 The Mössbauer data for all three complexes, 1, 2 and 3, at 77 K are shown in Table 5.The spectrum for 1 can be resolved into two quadrupole doublets, which implies the presence in the solid state of two non-equivalent high-spin iron(II) species. This may be a consequence of the disorder of the di(methylene) bridges in the tmen. Crystallographically, there must be centrosymmetry averaged over all the molecules in the unit cell.However, Fig. 1 indicates four possibilities: centrosymmetric containing C(2a) and C(2a9); centrosymmetric containing C(2b) and C(2b9); non-centrosymmetric containing C(2a) and C(2b9); and non-centrosymmetric containing C(2b) and C(2a9). Therefore half the iron centres have a centrosymmetric environment and half do not. This hypothesis explains the two species present in equal amounts suggested by the Mössbauer spectrum.The Mössbauer spectra obtained for 2 were also sharp and like that of 1 were resolved into pairs of ‘nested’ doublets. As with 1, the best fit is compatible with small differences in the symmetry of the d-electron cloud around the metal nuclei in each of the two solid-state components of the crystal. The relative areas of the two doublets suggested a rough 2 : 1 ratio of Fig. 5 Plot of the inverse of the atomic magnetic susceptibility (cA 21, m) and effective magnetic moment (meff, s) versus temperature for [FeCl5(tmen)3][BPh4] 3 (a) in the solid state and (b) in [2H8]thf solution.Diamagnetic correction: 1.990 × 1024 emu/atom FeII non-equivalent iron centres, but there is no obvious reason for such non-equivalence. Each unit cell of 2 contains four non-centrosymmetric dimers and two centrosymmetric ones, and the centrosymmetric molecule is disordered in the tmen bridges. There is a 2 : 1 ratio of the two types of molecule, but this cannot be the explanation of the non-equivalence of the Mössbauer parameters, because structural disorder does not imply differences in the electron density distribution at the metal.Consideration of packing in the solid may provide a more satisfactory answer. In the packing diagram (Fig. 6) the environments of Fe(1) and Fe(19) (in the centrosymmetric dimer) are centrosymmetric. This is because over the whole crystal there is a statistically symmetric distribution of the possible orientations of the disordered tmen ligands and each centrosymmetric dimer is symmetrically surrounded by noncentrosymmetric molecules, which are not disordered.However, there is no centre of symmetry in the Fe(2)/Fe(3) molecule (the non-centrosymmeric dimer) and the molecules surrounding this dimer are not related by any symmetry element. In addition, because of the disorder in the Fe(1)/Fe(19) molecule, the iron centres in the non-centrosymmetric dimer can experience various environments depending upon the orientations of the tmen bridges in the nearest centrosymmetric neighbours.Therefore, the environment of Fe(2) ? ? ?Fe(3) is not centrosymmetric. The 2 : 1 proportion holds here, because there are twice as many asymmetric surroundings as centrosymmetric ones. This could affect the symmetry of the electron density around each non-equivalent Fe nucleus, and hence the quadrupole splittings in the Mössbauer spectra. Fig. 6 Packing of the centrosymmetric (i) and non-centrosymmetric (ii) molecules of complex 2 (view down the a axis).The disordered methylene bridges and methyl groups of the chelating tmen molecules in the centrosymmetric dimer are depicted1986 J. Chem. Soc., Dalton Trans., 1997, Pages 1981–1988 We know of no precise analogy for the splitting of a Mössbauer signal due to asymmetric packing. However, there is at least one case in which iron atoms that might be expected to be similar do give rise to different signals.The complex [{Fe- (tmen)(O2CMe)}2(m-H2O)(m-O2CMe)2] is unexpectedly asymmetric in the crystal because the bridging water molecule hydrogen bonds more strongly to one iron-bound terminal carboxylate than to the other. The reason is not obvious, but the Mössbauer spectrum gives rise to two doublets, probably an electronic consequence of this asymmetry. The analogous complex with benzoate rather than of ethanoate is completely symmetrical in the crystal and gives rise to the expected single doublet in the Mössbauer spectrum.20 The Mössbauer spectrum of 3 was always a doublet, slightly broader than might have been expected of a single species.Attempts to fit the spectrum to a pair of ‘nested’ doublets were sometimes successful, although no satisfactory explanation was found for such behaviour. The lines were not sharpened by recrystallisation of the sample or by grinding the crystals with boron nitride and the broadening may be due to magnetic relaxation, which is common in high-spin FeIII compounds, though not in FeII.21 The isomer shifts and the quadrupole splittings of 1, 2 and 3 show the expected trends.The lower co-ordination number in the trigonal-bipyramidal 2 gives the lowest isomer shift because the s-character of the hybrid orbitals on the five-coordinate iron is greater. The low symmetry around iron reduces delocalisation of the d-electrons, consistent with the large quadrupole splittings.Five-co-ordinate environments are recognised to give rise to large quadrupole splittings.22 A combination of unique steric and electronic features makes tmen a very useful ligand for first-row transition metals such as vanadium and iron. Extensions into titanium(II),23 chromium(II),24 cobalt(II),7 and manganese(II) 7 chemistry have also been made. The iron(II)–tmen complexes show that solidstate adducts of distinct nuclearities can be isolated, though they are easily interconvertible in solution.These interconversions have equivalents in VII chemistry, but the thermal stability of the amine complexes seems to decrease from vanadium(II) to iron(II). A vanadium analogue of the dinuclear complex 2 has not yet been found, neither has the tetranuclear vanadium equivalent of [Fe4Cl8(thf)6]. Thus vanadium(II)–tmen chemistry still shows a number of big gaps. Nevertheless, the close parallels in the chemistry that we have already discovered suggest that these gaps might soon be filled, and that heteronuclear species containing both iron and vanadium should be accessible.We shall report shortly on our attempts to produce such materials. Experimental All operations were carried out under an inert atmosphere in a dinitrogen-filled drybox (Faircrest Engineering, Croydon) or with use of standard Schlenk techniques. Solvents were dried by standard procedures 25 and distilled under N2 prior to use. Iron(II) chloride was synthesised from metallic iron and HCl, and sodium tetraphenylborate (Aldrich) was used as received.N,N,N9,N9-Tetramethylethylenediamine was refluxed over molten sodium for ca. 1 h and then distilled under N2. Microanalyses were carried out at the Department of Chemistry, University of Surrey, using a Leeman CE 440 CHN elemental analyser. Chlorine contents were determined by Butterworth Laboratories (Teddington, Middlesex), while iron analyses were performed by Southern Science plc, using ICP– OES.Infrared data were recorded on a Perkin-Elmer 883 instrument, from Nujol mulls prepared under dinitrogen and spread on KBr plates. Conductivity measurements were carried out in tetrahydrofuran or dichloromethane solutions (ca. 1023 mol dm23) using a V-shaped cell (cell constant = 1.54) connected to a Portland electronic bridge. Mass spectra were recorded on a VG Autospec spectrometer (Fisons Instruments), equipped with a CsI gun at 25 kV (SIMS technique) or on a Kratos MS80RF machine with xenon at 8 kV (FAB technique).In both cases, 3-nitrobenzyl alcohol was used as the matrix. Mössbauer data were recorded at 77 K on an ES-Technology MS105 spectrometer with a 25 mCi 57Co source in a rhodium matrix. Spectra were referenced against iron foil at 298 K. Samples were pure solids or mixtures with boron nitride (ca. 50% w/w), ground to fine powders and then transferred to the aluminium sample holders in the glove box. The program used for spectra fitting and parameter calculation was ATMOSFIT 4.26 Magnetic susceptibility measurements were carried out in solution by variable-temperature NMR spectroscopy using the Evans method,27 with tetramethylsilane as the marker molecule unless otherwise stated.Samples were carefully weighed under N2, dissolved in the appropriate solvent mixture (deuteriated solvent and marker) in the glove box, and transferred to the outer tube of the coaxial NMR tube system. The solvent mixture without the paramagnetic sample was then placed in the inner tube.Spectra were recorded in the temperature range 100 to 50 8C (173–323 K) for [2H8]thf solutions. Measurements in the solid state for the same samples were carried out at the University of Surrey, using a Newport variable-temperature Gouy balance over the range 90–295 K. The field strength was calibrated by measurements on Hg[Co(NCS)4] and [Ni(en)3][S2O3] (en = ethylenediamine); the temperature scale was checked with CuSO4?5H2O.Corrections were applied for the diamagnetism of the sample tubes. Preparation of [FeCl2(tmen)2] 1 and [{FeCl(tmen)}2(Ï-Cl)2] 2 To a pink suspension of anhydrous FeCl2 (4.9 g, 38.5 mmol) in thf (180 cm3) heated under reflux, an excess of tmen (18 cm3, 13.9 g, 119.6 mmol) was added by syringe. The mixture changed to a light yellow solution. The heating under reflux was continued for 1.5 h, and the hot mixture was then filtered through Celite. A fine brown residue was separated from the light green filtrate and discarded. The solution was then concentrated under vacuum (at 30–40 8C) to ca. 120 cm3 and left standing at room temperature for 20 h. Very light green thin needles of the dinuclear complex 2 were the first isolated by filtration. They were washed with thf–hexane (1 : 2) and dried under vacuum. Yield: 1.6 g (Found for 2: C, 29.7; H, 6.90; N, 11.6. C12H32Cl4Fe2N4 requires: C, 29.7; H, 6.65; N, 11.5%). This crystalline product is highly oxygen- and moisture-sensitive, both in solution and in solid state, changing quickly to deep brown when exposed to air.After the isolation of 2, the filtrate was left at room temperature for a further 5 d, giving a mixture of light green needles and colourless thick prisms. One day later, only a homogeneous batch of colourless crystals of 1 was seen. They were then filtered off without washing or extensive drying. Additional amounts of 1, but not of 2, were obtained upon concentration of the mother liquor and cooling at 220 8C for a few days.Yield of 1 7.0 g (Found: C, 40.1; H, 9.30; N, 15.6. C12H32Cl3FeN4 requires: C, 40.1; H, 9.00; N, 15.6%). Total yield of both materials, based on the total content of FeII: 68%. IR (Nujol mull, selected absorptions) 1: n(C]N) at 1011s and 1026s, n(Fe]N) at 454m and 479m; 2: n(C]N) at 1009s and 1025s, n(Fe]N) at 441m, 462m and 485m cm21. FAB mass spectra for 2 (m/z, relative intensity): 449 [{M 2 Cl}1, 2], 269 (12); 207 [{FeCl(tmen)}1, 13]; 117 [{Htmen}1, (100)]; 72 [{Me2NCH2CH2}1, 58]; 58 [{Me2NCH2}1, 42%].Complex 2 was also prepared from an equimolar mixture of anhydrous FeCl2 (4.1 g, 32.3 mmol) and tmen (4.9 cm3) in refluxing thf (150 cm3). The crystals obtained were beautifulJ. Chem. Soc., Dalton Trans., 1997, Pages 1981–1988 1987 pale green long prisms which were washed and dried without any visible change. Yield: 6.3 g (80%). This procedure did not produce compound 1.Complexes 1 and 2 are soluble in thf, acetone and dichloromethane, insoluble in diethyl ether and hexane, and soluble with subsequent reaction in methanol. Neither compound is a conductor in solution in thf or CH2Cl2. Preparation of [Fe3Cl5(tmen)3][BPh4] 3 A colourless solution of Na[BPh4] (0.5 g, 1.5 mmol) in thf (10 cm3) was added slowly, at room temperature, to a very light green-yellowish solution of [FeCl2(tmen)2] 1 (1.5 g, 4.2 mmol) completely dissolved in thf (25 cm3).The reaction was carried out for ca. 20 h, giving a fine light brown suspension which was then concentrated to 25 cm3 under vacuum and filtered. The filtrate was carefully layered with 30 cm3 of hexane and left standing at room temperature for 7 d to produce long light brown prisms; they were isolated by filtration, washed with thf– hexane (1.5 : 1) and dried quickly to remove the excess of solvent. These prisms, like crystals of 1, become opaque and change to a powder when dried under vacuum.Yield: 0.9 g (64%). Purification of the product was achieved by recrystallisation from thf–hexane mixtures (Found: C, 50.1; H, 6.95; N, 8.20. C42H68BCl5Fe3N6 requires: C, 49.8; H, 6.80; N, 8.30%). Complex 3 was also synthesised in 62% yield from dinuclear 2 and Na[BPh4] under similar experimental conditions (Found: C, 49.7; H, 6.90; N, 8.40%). IR (Nujol mull): 1948w, 1878w, 1814w, 1764w; 1581ms; n(B]aryl) at 1429s and 1424s; n(C]N) at 1020s and 1004s; d(C]H aromatic, out-of-plane) at 749s, 735s and 708s; n(Fe]N) at 495m, 468m and 440m cm21.FAB mass spectrum (m/z, relative intensity): 553 (4); 269 (12); 154 (17); 136 (17); 117 [{Htmeda}1, 100]; 72 [{Me2NCH2- CH2}1, 90]; 58 [{Me2NCH2}1, 82%]. X-Ray diVraction analyses trans-[FeCl2(tmen)2] 1. Crystal data: C12H32Cl2FeN4, M = 359.2, monoclinic, space group P21/n (equivalent to number 14), a = 9.405(1), b = 12.333(1), c = 8.012(1) Å, b = 97.369(8)8, U = 921.7(2) Å3, Z = 2, Dc = 1.294 g cm23, F(000) = 384, m = 11.1 cm21, l(Mo-Ka) = 0.710 69 Å, T = 295 K.Crystals were air sensitive, clear, colourless plates with rounded edges. One, ca. 0.21 × 0.45 × 0.57 mm, covered by grease and mounted in a glass capillary tube containing a small amount of tmen, was examined photographically, then transferred to an Enraf-Nonius CAD-4 diffractometer (monochromated radiation). Accurate cell parameters were determined from the settings of 25 reflections q, in the range 10–118, and each reflection centred in four orientations.Diffraction intensities were recorded in the w–q scan mode to qmax = 258. During processing, corrections were applied for Lorentzpolarisation effects, absorption (by semiempirical y-scan methods) and to eliminate negative net intensities (by Bayesian statistical methods). There was no deterioration of the crystal. Of the 1621 unique reflections input into SHELX,28 1377 were ‘observed’ with I > 2sI.By analogy with the structure of trans- [VCl2(tmen)2],10 with which the crystals of 1 are closely isostructural, the Fe atom was placed at the origin and the remainder of the structure was located in successive difference maps. One end of the tmen ligand bridge is disordered equally in two orientations, but the N and the methyl-C atoms have not been resolved into separate sites for the two orientations. Hydrogen atoms have been included on all the methyl-C atoms (and refined with geometrical constraints) but not on the bridging methylene groups.The non-hydrogen atoms (except for the disordered bridging-C atom) were allowed anisotropic thermal parameters. At convergence, R = 0.052 and Rg 28 = 0.074 for all 1621 reflections weighted w = (sF 2 1 0.000 63F2)21. In the final difference map, the only peak of significance, at ca. 0.55 e Å23, is close to the bridge of the tmen ligand. For all the analyses, scattering factor curves for neutral atoms were taken from ref. 29. Computer programs used in this analysis have been noted above and in Table 4 of ref. 30, and were run on the MicroVAX 3600 machine in the Nitrogen Fixation Laboratory. [{FeCl(tmen)}2(Ï-Cl)2] 2. Crystal data: C12H32Cl4Fe2N4, M = 485.9, monoclinic, space group P21/n (equivalent to number 14), a = 9.9431(7), b = 23.152(2), c = 14.459(1) Å, b = 96.141(6)8, U = 3309.3(4) Å3, Z = 6, Dc = 1.463 g cm23, F(000) = 1512, m = 18.1 cm21, l(Mo-Ka) = 0.710 69 Å, T = 295 K. Crystals were air sensitive, pale green prisms.Several were sealed in glass capillaries in the glove-box. One, ca. 0.10 × 0.10 × 0.55 mm, was examined photographically then transferred to the CAD-4 diffractometer. Accurate cell parameters were determined as described above and diffraction intensities were recorded to qmax = 238. During processing, corrections were applied as above and for crystal deterioration (ca. 15% overall). Of 4585 unique reflections entered into the SHELX28 program system, 3094 had I > 2sI. The structure was determined by the automated Patterson routines in SHELXS 30 and refined by full-matrix least-squares methods in SHELXN.28b In one tmen ligand, there is disorder of a Me2NCH2 group equally in two orientations (with a common N atom site).Hydrogen atoms were included in all the tmen ligands except in the disordered group; methylene H atoms were in idealised positions, and those in methyl groups were refined with geometrical restraints.The isotropic temperature factors of all the H atoms were refined freely. All the non-hydrogen atoms, including the disordered C atoms, were allowed anisotropic thermal parameters. Refinement converged smoothly to R = 0.069 and Rg 28 = 0.076 for all 4585 reflections weighted w = (sF 2 1 0.001 51F2)21. In the final difference map, the highest peaks, at 0.8 e Å23 and <0.6 e Å23, were all within the tmen ligands. [Fe3Cl5(tmen)3][BPh4] 3. Crystal data: C18H48Cl5Fe3N6C24- H20B, M = 1012.7, monoclinic, space group P21/n (equivalent to number 14): a = 19.752(2), b = 21.617(2), c = 11.977(1) Å (cell dimensions from first crystal, see below), b = 99.339(8)8, U = 5045.7(9) Å3, Z = 4, Dc = 1.333 g cm23, F(000) = 2120, m = 11.5 cm21, l(Mo-Ka) = 0.710 69 Å, T = 295 K.Crystals were air sensitive, very pale yellow rectangular prisms; several were covered by grease and mounted in capillary tubes. After photographic examination, one, ca. 0.14 × 0.17 × 0.70 mm, was transferred to the CAD-4 diffractometer for the determination of accurate cell parameters as described above.Diffraction intensities were measured to qmax = 208, at which stage the intensities of three monitoring reflections had decreased by ca. 28.7%. For measurement of the data in the range 20 < q < 258, a second crystal, ca. 0.45 × 0.45 × 0.70 mm was used; this crystal deteriorated by ca. 8.0% during the data collection. The intensities for both crystals were corrected for Lorentz-polarisation effects and absorption (by semiempirical y-scan methods) before being scaled together, merged and adjusted by Bayesian statistical methods to ensure no negative net intensities. The combined data-set of 8871 unique reflections (5011 of which had I > 2sI) was then entered into the SHELX program system.28 The structure was determined from the direct methods routines in SHELXS31 and refined by large-block-matrix leastsquares methods with a total of 690 refined parameters.All non-hydrogen atoms were allowed anisotropic thermal parameters. Hydrogen atoms of methylene and phenyl groups were included in idealised positions and were set to ride on their parent carbon atoms; those in methyl groups were refined with1988 J. Chem. Soc., Dalton Trans., 1997, Pages 1981–1988 geometrical constraints. The isotropic temperature factors of all the hydrogen atoms were refined independently. At completion, R = 0.094 and Rg 28 = 0.077 for all 8871 reflections weighted w = (sF 2 1 0.0009F2)21.In a final difference map, the highest peaks, ca. 0.55 e Å23, were all close to the Fe3Cl5 core of the cation. Atomic coordinates, thermal parameters, and bond lengths and angles have been deposited at the Cambridge Crystallographic Data Centre (CCDC). See Instructions for Authors, J. Chem. Soc., Dalton Trans., 1997, Issue 1. Any request to the CCDC for this material should quote the full literature citation and the reference number 186/481.Acknowledgements We thank Mrs. J. E. Barclay and Dr. D. J. Evans for help with Mössbauer measurements and Ms. N. Walker for the microanalyses, and we are grateful for financial assistance (to J. S. de S.) from the British Council and the Brasilian Research Council (CNPq), and to the BBSRC for support. References 1 D. L. Hughes, L. F. Larkworthy, G. J. Leigh, C. J. McGarry, J. R. Sanders, G. W. Smith and J. S. de Souza, J. Chem. 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Sheldrick, Acta Crystallogr., Sect. A, 1990, 46, 467. Received 9th December 1996; Paper 6/08285H

ISSN:1477-9226    

DOI:10.1039/a608285h

出版商:RSC    

年代:1997

数据来源: RSC