Phosphine Ligands

Phosphine Ligands • Tertiary phosphines, PR3, are an important class of ligands because their electronic and steric properties can be altered in a systematic and predictable way over a very wide range by varying the R group(s). • Tertiary phosphines also stabilize an exceptionally wide variety of metal complexes of interest to the organometallic chemist as their phosphine complexes (R3P)nM−L. • Phosphines are more commonly spectator rather than actor ligands. 1 P P H H Me H phosphine MeO F trif luorophosphine P P P OPh OPh PhO tricyclohexylphosphine P P P Me Me Me Me Cy Cy Cy triphenylphosphite P Me Me methyldiphenylphosphine P OMe OMe trimethylphosphite Me Me Ph Ph Me triphenylphosphine P F P Ph Ph Ph trimethylphosphine P F P Me Me Ph Ph Ph Ph * Me dimethylphosphinomethane (dmpm) P dimethylphosphinoethane (dmpe) P Ph Ph P P Ph Ph PROPHOS P P Ph Ph Ph Ph Ph Ph * Me dimethylphosphinomethane di h l h hi h (dmpm) dimethylphosphinoethane di h l h hi h (dmpe) P Me Ph Cy Ph2P * * O DIPAMP PAMP PPh2 Ph-4-OMe Ph 4-MeO-Ph Ph Me S,S-chiraphos P P * Ph Ph * O Me Me DIOP R R R R *P P* *P P* R R R R 2 DuPHOS • The range and complexity of phosphine ligands has increased dramatically in recent years due to the development of C-X coupling reactions such as Buchwald-Hartwig reaction. 3 • Like NR3, phosphines have a lone pair on the central atom that can be donated to a metal. • Unlike NR3, they are also π-acids, to an extent that depends on the nature of the R groups present on the PR3 ligand. • For alkyl phosphines, the π acidity is weak; aryl, dialkylamino, and alkoxy groups are successively more effective in promoting π acidity. • In the extreme case of PF3, the π acidity becomes as great as that found for CO! In the case of CO the π* orbital accepts electrons from the metal. The σ* orbitals of the P−R bonds play the role of acceptor in PR3. 4 • Whenever the R group becomes more electronegative, the orbital that the R fragment uses to bond to phosphorus becomes more stable (lower in energy). • This implies that the σ* orbital of the P−R bond also becomes more stable. • At the same time, the phosphorus contribution to σ* orbital increases, and so the size of the σ* lobe that points toward the metal increases • Both of these factors make the empty σ* more accessible for back donation. • The final order of increasing π-acid character is PMe3 ≈ P(NR2)3 < PAr3 < P(OMe)3 < P(OAr)3 < PCl3 < PF3 ≈ CO 5 • The empty P−R σ* orbital plays the role of acceptor in metal complexes of PR3. • As the atom attached to the P atom becomes more electronegative, the empty P−X σ* orbital becomes more stable (lower in energy) making it a better acceptor of electron density from the metal centre. 6 • • • Occupation of the P−R σ* orbital by back donation from the metal also implies that the P−R bonds should lengthen slightly on binding. In practice, this is masked by a simultaneous shortening of the P−R bond due to donation of the P lone pair to the metal, and the consequent decrease in P(lone pair)–R(bonding pair) repulsions. LnM LnM-PR3 PR3 * * PR3 PR3 Once again, as in the case of CO, the M−L π bond is made at the expense of a bond in the ligand, but this time it is a σ, not a π, bond. 7 Tolman Electronic Parameter The electronic effect of various PR3 ligands can be adjusted by changing the R group as, quantified by Tolman, who compared the ν(CO) frequencies of a series of complexes of the type LNi(CO)3, containing different PR3 ligands. The increase in electron density at the nickel from PR3 σ-donation is dispersed through the M-L π system via π-backbonding. Much of the electron density is passed onto the CO π* and is reflected in decreased v(CO) stretching frequencies which corresponds to weaker CO bonds. CO stretching frequencies measured for Ni(CO)3L where L are PR3 ligands of different σ-donor abilities. [v(CO) =2143 cm-1] 8 Tolman Cone Angle • The second important feature of PR3 as a ligand is the variable steric size, which can be adjusted by changing R. • CO is so small that as many can bind as are needed to achieve 18e. In contrast, the same is rarely true for phosphines, where only a certain number of phosphines can fit around the metal. • This can be a great advantage in that by using bulky PR3 ligands, we can favor forming low-coordinate metals or we can leave room for small but weakly binding ligands, • The usual maximum number of phosphines that can bind to a single metal is
two for PCy3 or P(i-Pr)3
three or four for PPh3
four for PMe2Ph
five or six for PMe3 9 • Coordination Number (CN) – the number of bonding groups at metal centre
Low CN favored by: 1. Low oxidation state (e- rich) metals. 2. Large, bulky ligands. Although Pd(P(tBu)2Ph)2 is coordinatively unsaturated electronically, the steric bulk of both P(tBu)2Ph ligands prevents additional ligands from coordinating to the metal. 10 • The cone angle is obtained by taking a space-filling model of the M(PR3) group, folding back the R substituents as far as they will go, and measuring the angle of the cone that will just contain all of the ligand, when the apex of the cone is at the metal. • Although the procedure may look rather approximate, the angles obtained have been very successful in rationalizing the behavior of a wide variety of complexes. P M = Tolman's cone angle 11 Tolman Plot of Electronic Parameter and Cone Angle 12 • An important part of organometallic chemistry consists in varying the steric and electronic nature of the ligand environment of a complex to promote whatever properties are desired: activity or selectivity in homogeneous catalysis, reversible binding of a ligand, facile decomposition, or high stability. • Using the Tolman plot we can relatively easily change electronic effects without changing steric effects
e.g., by moving from PBu3 to P(OiPr)3 • Also, we can relatively easily change steric effects without changing electronic effects
e.g., by moving from PMe3 to P(o-tolyl)3 13 Metal Carbenes • The concept of a double bond between transition metals and carbon constitutes one of the most important elements in the field of organometallic chemistry • The notion of a metal–carbon double bond was first brought forward by Fischer and Maasbol in 1964 with the synthesis of (CO)5W=C(Ph)(OMe) • Soon after the discovery of Fischer type complexes their chemistry was systematically explored and they have been since well established as valuable species in organic synthesis as well as in catalytic processes • Schrock later prepared a number of tantalum complexes including (Np)3Ta=CH(CMe3) and (η5Cp)2MeTa=CH2 14 • Two different patterns of reactivity emerged during the development of these systems resulting in their classification as Fischer and Schrock type carbenes. • Each represents a different formulation of the bonding of the CR2 group to the metal and real cases fall somewhere between the two. Fischer type vs. Schrock type carbene complexes 15 OC OC CO CO O W CO CH3 W(0) , 18 VE CH3 CH3 Ta H H Ta(0) , 18 VE If we consider the Schrock carbene as a Fischer carbene with strongly enhanced M→ →C p-back bonding, the 2 electrons originally in M(dπ) transfer to the C(pz) orbital, oxidizing the metal by 2 units and giving a CR22- ligand. • The more electropositive early transition metal complexes have less stable M(dπ) orbitals, i.e. easier to oxidize. For example d2 metals are especially strong pdonors. • The system can thus be visualized as a metal stabilized carbanion acting as both a σ and π donor to the metal. 16 • Free carbene CH2 has two distinct spin isomers: singlet and triplet – not resonance forms (sinlget ↔ triplet resonance forbidden) • Singlet and triplet forms have different H-C-H angles • In the singlet state 2e- are paired up in the sp2 orbital leaving the pz orbital unoccupied • In the triplet state both the sp2 and p orbitals are singly occupied 17 Fischer • • vs. Schrock In the Fischer (singlet) case, direct C→M donation predominates and the carbon tends to be positively charged. In the Schrock (triplet) case, two covalent bonds are formed, each polarized toward the 18 carbon giving it a negative charge. • Taylor and Hall used ab-initio calculations to differentiate between the electronic structures of Fischer and Schrock type carbene complexes. • Calculations on a variety of free carbenes indicated that •
heteroatom and phenyl substituents preferentially stabilize a singlet ground state.
alkyl and hydride substituents stabilized a triplet ground state at the carbene carbon. Carbenes are both thermodynamically and kinetically unstable therefore forming very strong metal-carbene bonds disfavoring dissociation e.g. just as paramagnetic triplet :CH2 can dimerize to form diamagnetic H2C=CH2, it also binds to a triplet LnM fragment to give a diamagnetic LnM=CH2 complex. 19 Metal Carbynes • Have similar bonding formulations as per Fischer and Schrock carbenes • The free carbyne can be of doublet (Fischer) or quartet (Schrock) multiplicity • The carbyne ligand is linear • Carbyne carbon is sp hybridized • The M≡C bond is very short (1.65 – 1.90 Å) • Characteristic low-field 13C NMR resonance in the range +250 to +400 ppm 20 Fischer Carbynes • A doublet (Fischer) carbyne is sp hybridized • Contains one filled sp orbital capable of donating 2e- to a metal centre • Contains one singly occupied p orbital capable forming an additional π bond • The remaining empty p orbital is capable of M→C π back donation • 3e- donor covalent model (or 4e- donor ionic model) 21 Schrock Carbynes • A quartet (Schrock) carbyne is also sp hybridized • Contains three singly occupied orbitals (one sp and two p) capable of forming three covalent M-C bonds (one s and two p bonds) • 3e- ligand in covalent model (or 6e- ionic model) 22