1 13-1 Historical background The first organometallic compound reported was synthesized in 1827 by Zeise (asserted that this yellow product contained an ethylene group). – yellow needle-like crystals after refluxing a mixture of PtCl4 & PtCl2 in EtOH, followed by addition of KCl solution.
2 Bau et al, Inorg. Chem. 1975, 14, 2653. 3 Grignard Reagents Reactions between magnesium (Mg) & alky halides, performed by Barbier in 1898-9, & subsequently by Grignard, led to the synthesis of alkyl magnesium complexes now known as Grignard reagents (V. Grignard, 1912 Chemistry Nobel Laureate; P. Sabatier).
4 Sandwich compounds
5 Ferrocene (二茂鐵) Modern organometallic chemistry began in 1950s & 1960s., especially since the discovery of ferrocene by P. Pauson & S. A. Miller in 1951 & structurally characterized by E. O. Fischer (TUM, Germany; 1973 Chemistry Nobel Laureate; G. Wilkinson).
6 Projection from the Cn axis
Cyclo-C5H5MgBr + FeCl3 7 Cluster Compounds
8 More Examples of Organometallic Compounds
9 Organometallic Chemistry is a subfield of coordination chemistry in which the complexes contains an M-C or M-H bond.
Organometallic species tend to be more covalent, & the metal is often more reduced (low-valent), than in other coordination compounds.
10 Dorothy Hodgkin Vitamin B12 coenzyme : (1910-1994) naturally occurring & Nobel Prize in oldest cobalt complex, Chemistry (1964) contains acobalt-carbon σ bond.
It catalyzes 1,2-shift reaction
11 13-2 Organic ligands and nomenclature
12 - C3H5 (C3H6-H)
13 η : # of ligand atoms attached to the metal.
μ : # of metal atoms bridged. 14 13-3 The 18-electron rule
Main group chemistry : the octet rule (s2 + p6) Organometallic chemistry : 18-electron rule (s2 + p6 + d10)
CpFe(CH3)(CO)2 Neutral counting: Cp 5 + Fe 8 + CH3 1 + 2 CO 4 = 18
15 Electron Count for Common Ligands (ionic model is preferred) Neutral Ligands, L ⎯ (1) lone pair donors such as CO & NH3 (2) π-bond donors such as C2H4 (3) σ-bond donors such as H2 Anionic Ligands, X ⎯ Negatively charged ligands such as H-, Cl-, or Me-. For a M⎯X fragment, the ligand X is considered as a 2-e donor with a - 1 oxidation state, & the metal is considered as in a + 1 oxidation state (indeed, totally one-e donor).
16 -1
+1
NO NO - NO- NO NO+ NO+ 17 18 Method A Method B (Ionic model) (Covalent model)
- 4 C3H5
19 6 ♦ 2.1 contains a L3 type, η -benzene ligand; and 2.2 ~ 2.5 are associated with a LX type, η3-allyl ligand; complex 2.6 contains only an X type, η1-allyl ligand.
20 ♦ Bridging Ligands ⎯ (1) μ-L type, such as the bridging CO’s in (CO)3Co(μ-CO)2Co(CO)3, where the μ-CO is a 1-e donor to each metal. Other μ-L type ligands like μ-：CH2 in M⎯CH2⎯M′ is also a 1-e donor to each metal (metal-carbene, -carbyne).
21 (2) μ-LX type, such as the bridging Cl– in + LnM⎯Cl：→M ′Ln. Other μ-LX type ligands :
⎯OR, ⎯SR, or ⎯PR2 etc.
μ3 or μ4 ?
22 - - (3) μ-X1 type, such as the bridging H , or Me . For aM⎯X⎯M′ fragment is considered a 2-e, 3- centered bridge-type bonding.
(4) μ-X2 type, such as the bridging oxygen, M⎯O⎯M′, where the μ-O is a 2-e donor to each metal (totally a 4-e donor).
23 t2g
24 t2g
5 (η -C5H5)2Fe
25 26 + 5 - Charged species, [Mn(CO)6] , [(η -C5H5)Fe(CO)2] 18e 18e
(CO)5Mn⎯Mn(CO)5
+ ClMn(CO)5; [Re(CO)5(PF3)]
27 13-3-2 Why 18 electrons ? Main group chemistry : the octet rule (s2p6)
Organometallic chemistry : 18-electron rule (s2p6d10) This cannot provide an explanation for why so many complexes violate the 18-electron rule.
28 6 Cr(CO)6 (d / 18 e)
Lowering t2g orbitals
Δo
Raising eg orbitals
29 Cr(CO)6 : (1) The strong σ-donor ability of CO raises the
eg orbitals in energy (more antibonding) (2) The strong π-acceptor ability of CO lowers
the t2g orbitals in energy (more bonding)
Ligands that are both strong σ donors & π acceptors should therefore be the most effective at forcing adherence to the 18-electron rule.
30 (less bonding)
31 Exceptions :
2+ 1. [Zn(en)3] (22-electron species): en is a good σ donor, but not as strong as CO. As a result,
eg is not so antibonding to cause significant destabilization of the complex, & thus 22-e species is stable.
32 2- - 2. TiF6 (12-electron species) : F is a π-donor as well as a σ-donor. The π-donor ability of F-
destabilizes the t2g orbitals, making them slightly 2- antibonding. The species TiF6 has 12 electrons in the bonding σ orbitals and no electrons in the
antibonding t2g or eg orbitals.
π-donor
π-acceptor
33 The most common exception is square-planar geometry, in which a 16-electron configuration may be the most stable, especially for complexes of d8 metals.
34 Δsp antibonding
35 Consequently, for square-planar complexes of ligands having both σ-donor & π-acceptor characteristics, a 16-electron configuration is more stable than an 18-electron configuration.
It is also possible to accept one ligand to achieve an 18 (20)-electron configuration as well as a 5 (6)- coordinate configuration is formed.
36 13-4 Ligands in organometallic chemistry 13-4-1 Carbonyl (CO) complexes Carbon monoxide is the most common ligand in organometallic chemistry (both a σ-donor & a π-acceptor).
C O 37 Bonding : Two features of molecular orbitals needs attention about CO. 1. The HOMO has its largest lobe on carbon. 2. Carbon monoxide also has two empty π* orbitals (LUMO); these also have larger lobes on carbon than on oxygen.
HOMO LUMO
38 39 The bond strength of CO ? (1) X-ray crystallography (less sensitive) free C≡O 112.8 pm metal carbonyls 115 pm (2) Infrared spectrum (more sensitive)
40 ν(CO) = 2143 cm-1 for free C≡O
Why ??
Isoelectronic species (d6)
41 Actually, calculations have demonstrated that it is much more likely that donation from the HOMO to the metal in cationic complexes is insignificant in comparison with the polarization effect caused by the metal cations.
Polarization large Polarization small More covalent 42 The presence of a transition metal cation tends to reduce the polarization in the C≡O bond by attracting the bonding electrons. The consequence is that the electrons in the positively charged complex are more equally shared by the carbon & the oxygen, giving rise to a stronger bond & a higher-energy C≡Ostretch.
43 Bridging modes of CO The bridging mode is strongly correlated with the position of the CO stretching band.
In cases where CO bridges two metal atoms, both metals can contribute electron density into π* orbitals to weaken the CO bond & lower the energy of the stretching.
44 45 46 Electron count
47 Mo-Mo triple bond Mo-Mo single bond
Electron Count ? Mo-Mo Mo≣Mo Mo 6 6 Cp 5 5 CO 3x2 2x2 Mo ? Mo bond 1 3 Total 18 18
48 Binary carbonyl complexes
49 V(CO)6 (17-electron configuration)
In V(CO)6, it is apparently too small to permit a seventh coordination site; hence, no metal-metal bond dimer can occur to give an 18-electron configuration. However, V(CO)6 is easily reduced - to [V(CO)6] , a well-studied 18-electron example.
50 An interesting feature of the structures of binary carbonyl complexes is that the tendency of CO to bridge transition metals decreases in going down the periodic table.
I.e., in Fe2(CO)9 – 3 bridging CO in M2(CO)9 (M = Ru, Os) – 1 bridging CO A possible explanation is that the orbitals of bridging CO are less able to interact effectively with transition metal atoms as the size of the metals increases.
51 Synthesis of Binary Carbonyls
1. Direct reaction of a transition metal with CO.
Ni + 4 CO ' Ni(CO)4 (At RT & 1 atm)
Ni(CO)4 is a volatile, extremely toxic liquid. Because the reaction can be reversed at HT, coupling of the forward & reverse reactions has been used commercially in the Mond process for obtaining purified nickel from ores.
52 2. Reductive carbonylation.
CrCl3 + 6CO + Al → Cr(CO)6 + AlCl3
Re2O7 + 17CO → Re2(CO)10 + 7CO2 3. Thermal or photochemical reaction of other carbonyls.
53 Ligand substitution
An excellent precursor
(Mn)(bpy) (Mn) (bpy) A common reaction is replacement of the lost CO by another ligand to form a new 18-e species. This type of reaction therefore provides a pathway in which CO complexes can be used as precursors for a variety of complexes of other ligands.
54 Oxygen-bonded carbonyls This phenomenon was first noted in the ability of the oxygen of a metal-carbonyl complex to act as a donor toward Lewis acids, with the overall function of CO serving as a bridge between the two metals.
oxophilic metals 55 Attachment of a Lewis acid to the oxygen results in significant weakening & lengthening of the C-O bond & a corresponding shift of the C-O stretching vibration to lower energy in the infrared. This shift is typically between 100 & 200 cm-1.
weakened weakened
56 13-4-3 Hydrides and dihydrogen complexes Both ligands have played important roles in the development of applications of organometallic chemistry to organic synthesis (catalytic processes). Dihydrogen
Hydrides
57 Hydride complexes
58 One of the most interesting aspects of transition metal hydride chemistry is the relationship between this ligand & the rapidly developing chemistry of the dihydrogen ligand, H2.
oxidative addition (v.s. reductive elimination)
59 Dihydrogen complexes H-H = 0.84 Å (neutron) 0.74 Å (free) [0.82-0.90 Å for
coordinated H2]
Kubas et al, J. Am. Chem. Soc., 1984, 106, 451. 60 61 If the metal is electron rich & donates strongly to the σ* of H2, the H-H bond in the ligand can rupture, giving separate H atoms (hydrides).
The Kubas Complex, W(CO)3(P-i-Pr3)2(H2), which shows that hydrogen can bind to transition metal complexes, has been called the most important development in inorganic chemistry in the last 20 years. 62 Consequently, the search for stable H2 complexes has centered on metals likely to be relatively poor donors, such as those in high oxidation states or surrounded by ligands that function as strong electron acceptors (CO &NO).
II
II trans-[Os (H2)Cl(dppe)2]PF6 63 Dihydrogen complexes have frequently been suggested as possible intermediates in a variety of reactions (i.e., steps in some catalytic processes). As this ligand becomes more completely understood, the applications of its chemistry are likely to become extremely important.
64 13-4-4 Ligands having extended π systems Although it is relatively simple to describe pictorially how ligands such as CO & PPh3 bond to metals, explaining bonding between metals & organic ligands having extended π systems can be more complex.
65 CO, PPh3; (extended) π systems – bond to metals
Linear π systems
66 67 Cyclic π systems
68 69 13-5 Bonding between metal atoms & organic systems
π-ethylene complexes
π π∗
70 C=C distance in Zeise’s salt is 137.5 pm in free ethylene is 133.7 pm (1) donation of electron density to the metal in a σ -fashion reduces the π-bonding electron density within the ligand, weakening the C=C bond.
71 (2) The back-donation of electron density from the metal to the π* orbital of the ligand also reduces the C=C bond strength by populating the antibonding orbital. The net effect weakens & lengthens the C=C bond in the C2H4 ligand. IR : 1516 cm-1 in the Zeise’s salt 1623 cm-1 in free ethylene
72 π-allyl complexes
electron count 73 2 (π*)
1 (πn)
0 (π)
74 2-e donor 4-e donor
Na/Hg
Mn2(CO)10
Mn 5 Mn 5 CO 12 CO 10 1 3 η -C3H5 1 η -C3H5 3 75 13-5-2 Cyclic π systems Cyclopentadienyl (Cp) complexes 5 FeCl2 + 2NaC5H5 → (η -C5H5)2Fe + 2NaCl
2Na + 2C5H6 → 2NaC5H5 + H2 η1-, η3-, & η5-
Cp* = C5(CH3)5 Pentamethylcyclopentadiene is a better electron-donor than Cp, & is more difficult to be removed from metal centers. 76 In developing the group orbitals for a pair of C5H5 rings, we pair up molecular orbitals of the same energy & same number of nodes (symmetry).
See Fig. 13-22
dyz
M : s & dz2 pz 77 Furthermore, in each pairing there are two possible orientations of the ring molecular orbitals : one in which lobes of like sign are pointed toward each other, & one in which lobes of opposite sign are pointed toward each other.
78 79 80 Totally 10 group orbitals
81 Totally 9 metal orbitals Antibonding
Totally 19 molecular orbitals Non-bonding
Weakly bonding ∵4s (Fe) is involved. There are 18 electrons here, the 19th & 20th electrons go to antibonding orbitals. 82 Antibonding
Mainly nonbonding (slightly antibonding)
Weakly bonding
83 More stable (18 e) - unreactive
19 & 20 electron species are more reactive, tending
to form 18 electron products. 84 Ferrocene, however, is by no means chemically inert. It undergoes a variety of reactions, including many on the cyclopentadienyl rings. In general, electrophilic substitution reactions are much more rapid for ferrocene than for benzene, an indication of greater concentration of electron density in the rings of the sandwich compound (Cp-).
85 Co3+ (d6) + 6 x 2 Co+ (d8) + 4 + 6
electrophilic substitution reactions are much more rapid for ferrocene than for benzene 86 4n+2 rule for aromatic compounds
87
Complexes containing cyclopentadienyl and CO ligand
Figure 13.35 89 Molecular complexes containing other cyclic π-system
2- cyclooctatetraenyl ligand: [C8H8]
90 13-5-3 Fullerene complexes As immense π systems, fullerenes were early recognized as candidates to serve as ligands to transition metals.
91 These compounds fall into several structural types: (I) Adducts to the oxygens of osmium tetraoxide.
C60(OsO4)(4-t-butylpyridine)2 - The first pure fullerene derivative to be prepared.
92 93 (II) Complexes in which the fullerene itself behaves as a ligand.
2 5 2 Ex. Fe(CO)4(η -C60), Mo(η -C5H5)2(η -C60), 2 [(C6H5)3P]2Pt(η -C60)
94 As a ligand, C60 behaves primarily as an electron- dificient alkene (or arene) & bonds to metals in a dihapto fashion through a C-C bond at the fusion of two 6-membered rings. The d electron density of the metal can donate to an empty antibonding orbital of a fullerene result- ing in elongation of the C=C bond distance, & this pulls the two carbons involved slightly away from the C60 surface.
95 In some cases, more than one metal can be attached to a fullerene surface.
96 Hexa-platinum(II) C60 compound Some complexes of other fullerenes
97 2 2 2 5 5 μ3-η , η , η -C60 Fe(η -C5H5)(η -C70(CH3)3)
98 (III) Compounds containing encapsulated metals. These may contain one, two, or three metals inside the fullerene sphere.
Ex. UC60, LaC82, Sc2C74, Sc3C82
99 These complexes are structural examples of “cage” organometallic complexes in which the metal is completely surrounded by the fullerene. Typically, complexes containing encapsulated metals are prepared by laser-induced vapor phase reactions between carbon & metals.
100 These compounds contain central metal cations surrounded by a fulleride (a reduced fullerene). @ symbol designates encapsulation.
U@C60 contains U surrounded by C60 Sc3@C82 contains 3 Sc atoms surrounded by C82 This designation indicates structure only & does not include charges on ions that may occur.
101 102 H2-incorporated fullerene
This compound is stable at RT under air or vacuum, but start to release the incorporated H2 slowly at 160°C.
Komatsu et al, J. Am. Chem. Soc., 2003, 125, 7152. 103 (IV) Intercalation compounds of alkaline metals. These contain alkaline metal ions, occupying interstitial sites between fullerene clusters.
Ex. NaC60, RbC60, KC70, K3C60
104 Molecular orbital calculations show that t1u & t1g LUMO levels are low-lying (an acceptor).
While increasing reduction with alkaline metal atoms, the conductivity or superconductivity of AxC60 increases to a maximum for x~3.
Tc for A3C60 phase is 18 K in K3C60 & then up to a maximum of 33 K in RbCs2C60.
105 13-6 Complexes containing M-C, M=C, & M≡C bonds
106 13-6-1 Alkyls and related complexes Examples : Grignard reagents (with magnesium -alkyl bonds), & alkyl complexes with alkali metals (methyllithium) are well known.
107 Synthetic methods : 1. Reaction of a transition metal halide with organo- lithium, -magnesium, or -aluminum reagent.
ZrCl4 + 4PhCH2MgCl → Zr(CH2Ph)4 + 4 MgCl2 2. Reaction of a metal carbonyl anion with alkyl halide.
Na[Mn(CO)5] + CH3I → CH3Mn(CO)5 + NaI
108 Alkyl complexes have a tendency to be kinetically unstable & difficult to isolate; their stability is enhanced by structural crowding, which protects the coordination sites of the metal by blocking pathways to decomposition.
W(CH3)6 : stable up to 30 °C (6-coordinate) Ti(CH3)4 : decomposition at approximately -40 °C (4-coordinate)
109 Metallacycles are proposed as intermediates in a variety of catalytic processes.
110 111 13-6-2 Carbene complexes Carbene complexes contain metal-carbon double bonds. First synthesized in 1964 by E.O. Fischer, carbene complexes are now known for the majority of TMs & for a wide range of ligands, including the prototype carbene, :CH2.
Ernst Otto Fischer (1918-2007) The Nobel Prize in Chemistry 1973 (TUM)
112 The majority of such complexes, including those first synthesized by Fischer, contain one or two highly electronegative heteroatoms such as O, N, or S directly attached to the carbene carbon (Fischer-type carbenes).
X X X MC M C M C R R R And the presence of an electron withdrawing heteroatom on the carbene carbon makes the carbon slightly positive (Cδ+) & susceptible to nucleophilic attack. 113 π back bonding π-bond
σ-donor σ-bond
(L/2e lone pair donor) (X2/bis-alkyl)
empty Fischer carbene orbital empty filled orbital orbital filled orbital
R O C M C N
competition for π -backbonding from the metal and the lone pair orbital(s) on the functional group(s) to the carbene empty orbital (N and S the best, 114 then O, Ph, and other π -donating or lone pair containing groups) Schrock-type carbene complexes (alkylidenes): first synthesized several years after the initial Fischer carbene complexes by Schrock, contain only carbon and/or hydrogen attached to the carbene carbon. R sp2 C R' triplet state
formation of stable σ-bond & π-bond 115 δ+ δ-
116 Another aspect of bonding of importance to carbene complexes is that complexes having a highly electronegative atom such as O, N, or S attached to the carbene carbon tend to be more stable.
The stability of the complex is enhanced if the highly electronegative atom can participate in the π bonding, with the result a delocalized, 3-atom π system involving a d orbital on the metal & p orbitals on carbon & on the electronegative atom.
117 Such a delocalized 3-atom system provides more stability to the bonding π electron pair than would a simple metal-to-carbon π bond.
118 or CH3I
Fischer carbene Evidence for double bonding between Cr & C is provided by X-ray crystallography : Cr=C 204 pm, a typical Cr-C single bond distance of ca. 220 pm.
119 A single proton NMR at RT is as expected, but the splitting of this peak at LT into two peaks suggests two different proton environment due to the double-bond character of C-O.
120 A typical C=O double bond (116 pm) 121 13-6-3 Carbyne (alkylidyne) complexes Carbyne complexes have metal-carbon triple bonds; they are formally analogous to alkynes. M≡C-R Carbyne complexes were first synthesized fortuit- ously in 1973 by E. O. Fischer as products of the reactions of carbene complexes with Lewis acid. 122 carbyne 123 The best evidence for the carbyne nature of the complex is provided by X-ray crystallography, which gives a Cr-C bond distance of 168 pm (for X = Cl), considerably shorter than the 204 pm for the parent carbene complex (Cr≡C-C ~ 180° ). 124 The carbyne ligand has a lone pair of electrons in an sp hybrid on carbon; this lone pair can donate to form a σ bond. In addition, the carbon has two p orbitals that can accept electron density from d orbitals on Cr to form π bonds. Thus, the overall function of the carbyne ligand is as both a σ donor & π acceptor. 125 In some cases, molecules have been synthesized containing two or three of the types of ligands (alkyl, carbene, & carbyne). Such molecules may provide an opportunity to make direct comparison of lengths of metal-carbon single, double, & triple bonds. M-CR3 M=CR2 M≣CR 126 127 13-7 Spectral analysis & characterization of organometallic complexes Many complexes can be crystallized & characterized structurally by X-ray crystallography; however, not all organometallic complexes can be crystallized, & not all that crystallize lend themselves to structural solution by X-ray techniques. IR & NMR will be described in next section. 128 13-7-1 Infrared spectra IR can be useful in two respects : (i) The number of IR bands (molecular symmetry). (ii) The position of the IR band can indicate the function of a ligand & in the case of π-acceptor ligands, can describe the electron environment of the metal. 129 Linear or bent geometry? A single band indicates linear orientation of the CO ligands, & two bands indicate nonlinear orientation. 130 Complexes containing three or more carbonyls Although we can predict the number of IR-active bands by group theory, fewer bands may sometimes be observed (Table 13-7). 131 See Chapter 4 132 133 (I) Bands may overlap to such a degree as to be indistinguishable; (II) One or more bands may be of very low intensity & not really observed. (III) Isomers may be present in the same sample, & it may be difficult to determine which IR absorptions belong to which compound. 134 Positions of IR bands (2 bands) EN/IR stretching 135 The greater the electron density on the metal (& the greater the negative charge), the greater the back bonding to CO & the lower the energy of the carbonyl stretching vibrations. In combination with information on the number of IR bands, the positions of such bands for CO & other ligands can therefore be extremely useful in characterizing organometallic compounds. 136 137 138 2 : 2 : 1 139 140 141 Linear : 1 ; bent geometry : 2 142 P(C H ) 4 bands 6 5 3 143 P(C6H5)3 144 145 See Table 13-7 two mer (3) or fac (2) 146