Organometallic hemistry and omogeneous atalysis Dr. Alexey Zazybin Lecture N2 Kashiwa ampus, October 16, 2009 Properties I. Nature of II. Nature of L II.a. Ligands with -atom attached to II.b. Ligands with other atoms attached to metal
Alkylidene (carbene) ligand A free carben R 2 has two spin states (not resonance forms): singlet ( ) and triplet ( ) 2 R R sp 2 R sp 2 R pz pz Two types of coordinated carbene can be distinguished: Fischer type carbene and Schrock type carbene R R sp 2 R sp 2 R pz pz
Fischer s group, 1964 (O) 5 W O 3 W(O) 6 + 3 Li 3 ( 3 ) 3 O + (O) 5 W O 3 (O) 5 W O 3 3 Structural features of Fischer type carbenes: (O) 5 r=(oet)[n(i-pr) 2 ] 5 2 2.13 Å (r-r single bond distances are 2.0-2.2 Å) r OEt N(iPr) 2 1.35 Å (normal distance should be 1.41 Å, a 0.06 Å shortening) 1.33 Å (normal distance should be 1.45 Å, a 0.12 Å shortening) r OEt r OEt N(iPr) 2 N(iPr) 2
Fischer type carbene: 4 - middle to late transition metals - low oxidation state - π acceptor ligands at (O, NO) - π donor substituents (Oe, Ne 2 ) δ δ + R R Schrock, 1973 (t-butyl- 2 ) 3 Ta 2 + 2Li( 2 -t-butyl) Ta( 2 -t-butyl) 5 unstable intermediate α-elimination (t-butyl- 2 ) 3 Ta e e e + neopentane
Structural features of Schrock type carbenes: 5 p 2 Ta(= 2 )( 3 ) 138 Ta 3 2 2.24 Å The Ta= 2 bond is distinctly shorter than the Ta- 3 single bond! 2.03 Å Schrock type carbene: - early transition metals - high oxidation state - ligands with strong σ, π donor capacity - substituents with no π acceptor capacity (, R 3, SiR 3 ) δ + δ R R
Let s compare carbenes of Fischer and Schrock type: 6 Fischer arbenes Nucleophillic attacks at carbon atom of carbene (carbon is electron deficient) Electrophillic attacks on metal center (metal is more electron-rich, often d 6 18 e - system) Schrock Alkylidenes Electrophillic attacks at carbon atom of alkylidene (carbon is electron-rich) Nucleophillic attacks on metal center (metal is electron-deficient, usually d 2 or d 0 16 or 14 e - count) arbene is stabilized by heteroatom groups that can π- bond to it. Likes NR 2, SR, OR, or Ph groups. Later transition metals favored, especially with d 6 counts (carbene as neutral 2e - donor ligand) Alkylidene is destabilized by heteroatom groups that can π- bond to it. Strongly prefers or simple alkyl groups. Early transition metals favored, especially with d 0 centers (alkylidene as dianionic 4e - donor)
NR data for carbenes: 13 resonance for carbenes occurs to low field: 200-350 ppm, 7 1 resonance for =R occurs at +10 +20 ppm ompound 13, ass δ (ppm) p 2 Ta(= 2 )(e) 224 Schrock (t-bu 2 ) 3 Ta(=(t-Bu) 250 Schrock (O) 5 r(=(ne 2 )) 246 Fischer (O) 5 r(=ph(oe)) 351 Fischer (O) 5 r(=ph 2 ) 399 Fischer IR data: carbene complexes are characterized usually only through ν IR absorption (ν = is seldom unambiguously identified) o= 2 3080 and 2945 cm 1
Olefinic ligands 8 First organometallic compound Zeise salt, 1831 Pt K[Pt 3 ( 2 = 2 )]. 2 O Alkenes act as neutral 2e - donors (per = double bond). Due to the presence of empty π* antibonding orbitals, there is the possibility of some π-backbonding: d c-c, nm Dewar-hatt- Duncanson bonding model (1953) 2 2 K[Pt 3 ( 2 = 2 )] p* 2 Ti( 2 = 2 ) 1.337 1.375 1.438 σ-donation via the filled alkene π-system π-back donation via the empty alkene π -system 3 3 1.540
IR-spectroscopy in analysis of olefinic complexes Electronic effects with alkenes can often be s o m e what e a s i l y m o n i t o r e d using I R - spectroscopy. The more π-backbonding, the weaker the = double bond and the lower the = stretching frequency in the IR-spectrum Ethylene omplex ν = (cm 1 ) Free Ethylene 1623 [Ag(2=2)2]+ 1584 Fe(O)4(2=2) 1551 [Re(O)4(2=2)2]+ 1539 [pfe(o)2(2=2)]+ 1527 9 Pd24(2=2)2 1525 The table shows a series of alkene compounds with differing amounts of σ-donation and π - b a c k b o n d i n g. [Pt3(2=2)] 1516 pn(o)2(2=2) 1508 Pt24(2=2)2 1506 prh(2=2)2 1493 Factors influence thermodynamic stability of olefinic complexes: 1) Electron-withdrawing groups on the alkene generally increase the strength of the metal-alkene bonding 2) In cases where cis-trans isomerism is possible, the more stable complex is almost always formed by the cis-alkene (steric factor) 3) Third-row metals form the strongest bonds and most stable complexes
Arene ligands 0 Arenes typically coordinate in an η 6 fashion and as such are neutral 6 e- donors, although they can adopt lower coordination modes (η 4 and η 2 ). An interesting aspect of metal-arene complexes is that π-back donation plays a relatively important role in the bonding and chemistry.
Other types of coordination: 1 The reactivity of arenes in metals complexes differs from the free state: + 3 O long time heating + 3 O r(o) 3 reflux, 1-2 h - O 3 r(o) 3 Electron-withdrawing r(o) 3 or n(o) 3 fragments facilitate nucleophilic addition and substitution and other processes.
Sandwich bis-arene complexes 2 Bis-arene complexes are close to bis-cyclopentadiene complexes, but arene ligand is less tightly bonded to the metal center. r- 6 6 bond energy is 40 kcal/mol 6 6 5 5 Fe- 5 5 bond energy is 52 kcal/mol Ansa-metallocenes and metallocenophanes: When two rings (p or arene) are bridged together we have new class of organometallic compounds ansametallocenes. They have an application in the synthesis of stereoregular polymers, treating cancer and so on. r
arbonyl ligand Ludwig ond, 1888 3 Examples of complexes: Fe(O) 5, Ni(O) 4, o 2 (O) 8, n 2 (O) 10, r(o) 6 empty π*-acceptor orbitals on carbonyl powerful π-acceptor ligand! O excellent ligand, therefore, for stabilizing electron-rich low-valent metal centers Standard Bonding odes: O terminal mode 2e neutral donor O µ 2 bridging mode 2e neutral donor O µ 3 bridging mode 3e neutral donor
arbonyl IR Stretching Frequencies 4 The position of the carbonyl bands in the IR depends mainly on the bonding mode of the O (terminal, bridging) and the amount of electron density on the metal being π-backbonded to the O. The number (and intensity) of the carbonyl bands one observes depends on the number of O ligands present and the symmetry of the metal complex. Bonding odes: As one goes from a terminal O-bonding mode to µ 2 -bridging and finally µ 3 -bridging, there is a relatively dramatic drop in the O stretching frequency seen in the IR. O free O O terminal mode O O µ 2 bridging µ 3 bridging ν O IR (cm -1 ) 2143 2120-1850 1850-1720 1730-1500 (for neutral metal complexes)
The carbonyl region in the IR spectrum can be very distinctive and useful for help in assigning structures and for indicating the relative amount of electron density present on the metal: 5 Ni 2( µ -O)(O) 2(dppm) 2 O Ni O Ni O -O +O P P P P Ni (O) (dppm) 2 4 2 O O Ni P P P P Ni O O -O +O 1 2Ni(O) ( η -dppm) 3 O O Ni P P O 2000 1900 1800 Wavenumbers (cm-1) 1700
arbonyl NR Shifts 6 Terminal carbonyl ligands generally gives rise to 13 NR resonances in the region 160-230 ppm, the chemical shift (δ) being most significantly dependent on the metal center. In general, heavier metals cause shifts to the higher field. Bridging carbonyls typically resonate to lower field. O O O O 13 NR shifts: = Fe δ = 211 ppm = Ru = Os δ = 200 ppm δ = 183 ppm O Bonding -O could be drawn as a set of resonance forms: - O + O + O - X-Ray: d Ir- = 2.02 Å d Ir- = 1.87 Å [Ir(O) 5 ] 2+ d -O = 1.08 Å [(O) 2 Ir(S)(PPh 3 ) 2 ] + d -S = 1.51 Å
II.b. Ligands with other atoms attached to metal (, halogens, O, phosphines) 7 Phosphorus(III) ligands Examples: Phosphines: Pe 3, PPh 3 Phosphites: P(Oe) 3, P(OPh) 3 2 2 Ph 2 P PPh 2 -P bond formation is close to -(O): R R R P Phosphine ligands excellent soft-donor ligands with a wide variety of easily adjusted steric and electronic factors neutral 2e donor But in compare with O phosphines have more donor ability: Pe 3 > PPh 3 > P(Oe) 3 > P(OPh) 3 > P(NR 2 ) 3 > O ~ PF 3
Tolman, 1977 8 cone angle θ Phosphine/Phosphite one angel Ph 3 P PPh 3 Pd P 3 87 0 PF 3 104 0 P(Oe) 3 107 0 PPh 3 Ph 3 P PPh 3 Pd Ph 3 P PPh 3 Pe 3 118 0 PPh 3 145 0 P Pd P P(mesytil) 3 212 0
Structural features of phospine metal complexes 9 Some typical first row -PR3 average bond distances: Ti-P V-P r-p Ni-P 2.6 Å 2.5 Å 2.4 Å 2.1 Å NR 31 P characteristics of phosphine complexes 31 P NR shifts: + 245-163 ppm Downfield Region Upfield Region PPh (Oe) 2 PO 3 4 Pe Et 2 P 3 Pe 2 Pe 2 PPh 3 P 3 P(Oe) 3 O=Pe 3 PEt 3 Pe 3 P(N) 3 P 2 250 200 150 100 50 0-50 -100-150 -200-250 -300
ydride ligand Walter ieber,1931 O O O Fe O 0 ydride complexes quite a common case, many of hydride complexes are stable at ambient conditions - - hydride ligand anionic 2e donor Depending on the T, the nature of this ligand can vary from acidic (o(o) 4 is a strong acid) to hydride (like in Na): δ Μ Η δ+ δ+ Μ Η δ protic hydridic 1 -NR scale: δ, ppm +10 0-10
NR characterization of hydride complexes: ost hydride ligands show 1 NR signal in hydride region from -3 to -25 ppm, although it could be find in a wider range from +25 to -60 ppm. But it could not be a valid criteria of the acidity of metal-hydrogen bond 1 o(o)4 1 NR δ = 10.7 ppm p*2zr2 1 NR δ = + 7.5 ppm o(o) 4 + + o(o) 4 - p * 2Zr 2 + 2 3 I p * Zr 2 + 2 4 1 -NR scale: δ, ppm +10 0-10 oupling constants hydrides can couple with the atom, where this has ½ spin: 1) with other hydride ligand (if they are not equivalent) 2 J =1-10 z 2) with phosphorus atom (phosphines, phosphites) P 2 J P =15-30 z for cis-, 90-150 z for trans-orientation 3) with metal atom (Pt, Rh, W) P 1 J Pt =500-1500 z
IR characterization and types of hydride complexes 2 Terminal hydrides : o(o) 4 2- Fe 4 2- Re 9 Bridged hydrides: (O) 5 r r(o) 5 µ 2 hydride o o o o o o µ 6 hydride in [o 6 (O) 15 ] - - IR Spectra: 2200-1500 cm-1 can be weak or absent 1600-800 cm -1 broader (weak or absent) D-test: ν = 2 ν D Rep 2 ν Re =2000 cm -1 DRep 2 ν ReD =1460 cm -1
Structural Features (X-ray, neutron diffraction) 3 ydride is the smallest ligand and as a result, - distances are typically quite short: 1.8 to about 1.5 Å, depending on the metal. ydrides can be quite difficult to observe via X-ray diffraction (the most common technique used to determine structures) due to the very small number of electrons on the hydride vs. adjacent atoms, especially the metal. Since X-rays are scattered by electron density, not by atomic nuclei it is the - bonding electrons that are detected, so that X-ray methods systematically underestimate the true - internuclear distance by approximately 0.1 Å Therefore, neutron diffraction studies are considered best for accurately locating and identifying hydrides on metal centers (proton itself scatters neutrons relatively efficiently), even so much larger crystals are usually needed for neutron work (1 mm 3 vs. 0.01 m 3 )
Reactivity of hydride complexes 4
alogen ligands (halide donors) 5 The halides are anionic donors that generally only donate 2e - to a metal center. Due to their relatively high electronegativity they are not especially good σ-donor ligands. Although they can theoretically act as π-donor ligands, once again, the higher electronegativity limits them to simple 2e - donor ligands. Fluoride is generally NOT a good ligand except for very high oxidation state metal centers. It is too electronegative to donate much of its electron density. Bonding modes:
When coordinating to a single metal center usually we regard halides as 2e - donors. 4e - can be more easily donated by I - then by any other halide, even so exceptions are possible: 6 Bridging: oordinatively unsaturated metal halides often exist as insoluble polymers connected by bridging halides. Pd 2 is a classic example. Ligands such as 3 N can displace the bridging halide to give the soluble square planar complex ( 3 N) 2 Pd 2.
Oxygen donors: 7 Alkoxides commonly act as bridging ligands for electropositive metals. To avoid formation of bridges, bulky alkoxides must be used. With high valent transition metals, π-donation becomes important.
Important information about bonding type gives -O- angle: 8 Usually in ethers (like Et2O) the -O- angle is about 110 0 But in the metal-alkoxyde systems -O- angle is often approaching 120 or even 180 An angle near 120 suggests that the O is sp 2 hybridized and acting as a 4 electron donor. An angle near 180 suggests that the O is sp hybridized and acting as a 6 electron donor. In this case the alkoxide can be considered to be isoelectronic with the p ligand.
9
For organic compounds: maximum 8 electrons on : 0 s p For metal complexes: maximum 18 electrons on : s d p omplexes with 18 e - counts are referred to as saturated, because there are no empty low-lying orbitals to which another incoming ligand can coordinate. omplexes with less than 18e - are called unsaturated and can electronically bind additional ligands
ationic 2e- donor: NO + (nitrosyl) 1 Neutral 2e- donors: PR3 (phosphines), O (carbonyl), R2=R2 (alkenes), R R (alkynes, can also donate 4 e - ), N R (nitriles) Anionic 2e - donors: - (chloride), Br - (bromide), I - (iodide), 3 - (methyl), R3 - (alkyl), Ph - (phenyl), - (hydride) The following can also donate 4 e - if needed, but initially count them as 2e - donors (unless they are acting as bridging ligands): OR - (alkoxide), SR - (thiolate), NR2 - (amide), PR2 - (phosphide) Anionic 4e - donors: 35 - (allyl), O 2- (oxide), S2 - (sulfide), NR2 - (imide), R2 2- (alkylidene) and from the previous list: OR - (alkoxide), SR - (thiolate), NR2 - (amide), PR2 - Anionic 6e - donors: p - (cyclopentadienyl), O2 - (oxide)
2 Electron ount Oxidation State oordination Number Basic tools for understanding structure and reactivity. Doing them should be automatic. Not always unambiguous don t just follow the rules, understand them!
3 The basis of counting electrons Every element has a certain number of valence orbitals: 1 (1s) for 4 (ns, 3 np) for main group elements 9 (ns, 3 np, 5 (n-1)d) for transition metals s px py pz dxy dxz dyz dx2-y2 dz2
4 The basis of counting electrons Every orbital wants to be used", i.e. contribute to binding an electron pair. Therefore, every element wants to be surrounded by 2/8/18 electrons. The strength of the preference for electron-precise structures depends on the position of the element in the periodic table.
5 The basis of counting electrons Too few electrons: An empty orbital makes the compound very electrophilic, i.e. susceptible to attack by nucleophiles. Too many electrons: There are fewer covalent bonds than one would think (not enough orbitals available). An ionic model is required to explain part of the bonding. The "extra" bonds are relatively weak. etal-centered (unshared) electron pairs: etal orbitals are fairly high in energy. A metal atom with a lone pair is a strong σ-donor (nucleophile) and susceptible to electrophilic attack.
Use a localized (valence-bond) model to count electrons 6 2 Every has 2 e. OK 4 has 2 e, 8. OK N 3 N has 8 e. Nucleophile! OK N
2 4 has 8 e. OK 7 singlet 2 has only 6 e, and an empty p z orbital: extremely reactive ("singlet carbene"). Unstable. Sensitive to nucleophiles and electrophiles. triplet 2 has only 6 e, is a "biradical" and extremely reactive ("triplet carbene"), but not especially for nucleophiles or electrophiles.
3 + has only 6 e, and an empty p z orbital: extremely reactive. Unstable. Sensitive to nucleophiles. 8 3 - has 8 e, but a lone pair. Sensitive to electrophiles. - has 8 e, 4 lone pairs. OK Somewhat sensitive to electrophiles.
9 B 3 B has only 6 e, not stable as monomer, B forms B 2 6 : B 2 6 B has 8 e, all 's 2 (including the bridging Al 3!). 2-electron-3-center bonds! OK Al has only 6 e, not stable as monomer, forms Al 2 6 : B Al B Al 2 6 Al has 8 e, all 's too (including the bridging!). Regular 2-electron-2-center bonds! OK Al Al
2 eal 2 e 2 Al 2 4 0 e e Al Al Al Al e 2-electron-3-center bonds are a stopgap! 3 B N 3 e B N N-B: donor-acceptor bond (nucleophile N 3 has attacked electrophile B 3 ). Organometallic chemists are "sloppy" and write: B B Writing or would be more correct (although the latter does not reflect the real charge distribution). N B N N
1 P 5 P would have 10 e, but only has 4 valence orbitals, so it cannot form more than 4 net P- bonds. You can describe the bonding using ionic structures (hyperconjugation). Easy dissociation in P 3 en 2. P P? F 2 - Write as F F -, mainly ion-dipole interaction. F F? F F
Examples of electron counting: 2 Pt2(PPh3)2 1. Nature of central atom (oxidation state, number of electrons) harge of the complex: 0 harge of the nonneutral ligands: 2 * (-1) = -2 etal oxidation state: +2 Number of electrons for Pt 2+ : for Pt 0 according to the periodic table 10, For Pt 2+ : 8e - 2. Nature of ligands (number of electrons donating to the metal) 4 * (2e - - ligands) = 8e - Final result: 8 + 8 = 16 e - Try yourself: Ne 2 e 2 N o Ne 2 Ne 2 Ph 3 P Ph 3 P PPh 3 Ir F F Rh F F PhP P Ph 2 PPh Fe P Ph 2 Br +
3 Try yourself: Ne 2 e 2 N o Ne 2 Ne 2 Ph 3 P Ph 3 P PPh 3 Ir F F Rh F F PhP P Ph 2 PPh Fe P Ph 2 Br +
4 Exercises Give electron count and oxidation state for the following compounds. Draw conclusions about their (in)stability. e 2 g Pd(Pe 3 ) 4 ereo 3 Zn 4 Pd(Pe 3 ) 3 Zne 4 2- o(o) 4 - n(o) 5 - r(o) 6 V(O) 6 - V(O) 6 Zr(O) 6 4+ Pd(Pe 3 ) 3 Rh 2 (Pe 3 ) 2 OsO 4 (pyridine) Ni(Pe 3 ) 4 Ni(Pe 3 ) 3 Ni(Pe 3 ) 2 2