CHEM2009: Transition Metal ChemistryTopic 2: Transition Metal Chemistry and

an Introduction to Organometallic Chemistry

Professor Gill Reid

Lecture synopsis:

L1: History & relevance of organometallic chemistry; e-counting rules

L2&3: TM carbonyls, carbonylates & hydrides – preparation, bonding & spectroscopic features (NMR & IR spectroscopy); comparison of -acceptor ligands (CO, N2, RNC, NO); IR spectroscopy as a probe of ligand pr0perties

L4&5: TM alkyls, alkenes, alkynes, arenes – bonding properties, e-counting; case studies - examples

L6&7: Reaction mechanisms in TM chemistry; kinetics vs themodynamics; substitution reactions (A, D, I mechanisms); rates of water exchange & CFAE

L8: Substitution in organometallic complexes

L9: Oxidative addition & reductive elimination reactions; consequences & relevance in organometallic chemistry

L10: Migratory insertion and elimination (de-insertion) reactions – CO insertion into metal-alkyl bonds (or alkyl migration to M-CO bond)

L11: Homogeneous catalysis – combining elementary steps for industrial production of acetic acid via the Monsanto Process (BP)

Reading:• “Inorganic Chemistry” – Shriver & Atkins, Oxford, 5th Ed, 2009, Ch

21 & 22

• “Advanced Inorganic Chemistry” – F. A. Cotton, G. Wilkinson, C. A. Murillo and M. Bochmann, Wiley, 6th Ed,1999, aspects of Ch 16, 21, 22 and survey chapters

• “Organometallics I” – M. Bochmann, OCP 12, 1994

• “The Organometallic Hypertext Book” - http://www.ilpi.com/organomet/index.html

Lectures are supported by problem solving workshops and lab practicals which reinforce the key learning objectives and develop important skills in interpretation of various spectroscopic and other experimental data to provide logical rationale for the identification of the products from organometallic reactions.

These activities are fully integrated within the course, full participation is a requirement and all the relevant material is examinable.

History of organometallic chemistry & electron counting rules

L1

History of Organometallic Chemistry:

• 1760 1st organometallic compound (Cadet) [(CH3)2As]O

• 1827 1st alkene complex (Zeise) Na[PtCl3(C2H4)]

• 1849 1st alkyl Zn and Hg compounds (Frankland) [(C2H5)2Zn], [C2H5ZnI], [(CH3)2Hg]

• 1868 1st metal carbonyl complex (Schützenberger) [Pt(CO)Cl2]2

• 1890 1st binary metal carbonyl (Mond) [Ni(CO)4]

• 1909 1st -organotransition-metal compound (Pope) [(CH3)3PtI]

• 1951 1st sandwich complex (Pauson, Miller) [(C5H5)2Fe]

Nobel Prizes:

• 1963 Polymerisation catalysts (Ziegler, Natta)

• 1973 Sandwich complexes (Fischer, Wilkinson)

• 2001 Chiral catalysts (Knowles, Noyori, Sharpless)

• 2005 Metathesis catalysis (Chauvin, Grubbs, Schrock)

Effective Atomic Number Rule (EAN):

• Also called the Inert gas rule or the 18 electron rule

• “Thermodynamically stable transition metal organometallics are formed when the sum of the metal d electrons + the electrons conventionally regarded as being supplied by the ligands = 18"

• i.e. organometallic complexes are stable when the metal has effectively the atomic number of the next inert gas.

• [Similar to the Lewis octet rule for light main group elements]

[Only useful for organometallic compounds – do not attempt to use this in coordination chemistry]

Electron counting: ligands

• Consider ligands as Lewis bases: i.e. electron pair donors

– e.g. H will achieve an inert gas configuration as H- (hydride)

- a 2 electron donor (1s2)

– Cl will achieve a Lewis octet (inert gas config.) as Cl- (chloride)

- a 2 electron donor (3s23p6; 3pz = -donor)

– CH3 will achieve a Lewis octet (inert gas config.) as CH3- (isoelectronic with NH: CH3

- NH3 )

- a 2 electron donor (sp3 hybrids - with lone pair)

– CO is stable as neutral molecule (:C-O+)

- a 2 electron donor (sp hybrids - lone pair)

• Termed MLX approach – L = neutral 2e donor; X = charged 2e donor

Electron counting: complexes

• E.g.1 [Cr(CO)6]

– Ligands:

– 6 x neutral CO ligands - each 2 electron -donors

6 x 2 e count

– Metal: Cr(0)

– Cr(g) 3d54s1 - in complexes, d-s orbital separation is increased so we consider Co(0) in complexes as 3d6

6 e count

– Complex:

– Ligand + metal electron count = 18 e total

Electron counting, MLX approach

E.g.2 [IrCl2(CH3)(CO)(PPh3)2]

– Ligands:

• 3 anionic ligands - 2 Cl-, 1 CH3- - each 2 e -donors

• 3 neutral ligands (CO + 2 PPh3 - each 2 electron -donors)

12 e count

– Metal: Ir(III)

• Ir(g) 5d76s2 - in complexes, so lose 3 electrons and consider 5d6.

6 e count

– Complex:

• Ligand + metal electron count: 18 e total

Ligand nomenclature• ηn (Greek ‘eta’) number of ligand atoms, n, coordinated to a

metal;

– when n is absent, all unsaturated atoms in the ligand are assumed to be coordinated to the metal

– E.g. [Cr(η6-C6H6)2] i.e. dibenzenechromium

• μm (Greek ‘mu’) number of metal atoms, m, the ligand is coordinated to;

– when m is absent, the ligand is assumed to be coordinated to 2 metal atoms

E.g. [Co2(CO)8] i.e. di-μ-carbonyl-bis{(tricarbonyl)cobalt}

Electron counting: M-M bonds• E.g.3 [Rh(C5H5)(CO)2]

– Ligands:

• 1 anionic ligand (6 e donor)

• 2 neutral ligands (CO) 10 e count

– Metal: Rh(I)

• Rh(g) 4d85s1 - in complexes, and loses 1 electron to get +1; consider as 4d8. 8 e count i.e. 18 e overall

• E.g.4: [Rh2(C5H5)2(CO)3]

– Each metal has Rh(C5H5)(CO) unit (16 electrons).

– Bridging (μ2) CO donates an electron pair – 1 electron per metal: 17 electrons per unit (therefore unstable)

Have a 2-centre-2-electron Rh-Rh bond to achieve the 18 electron count

Why 18 electrons?

• Organometallic compounds usually involve strong-field ligands – i.e. good σ-donor or σ-donor/-acceptor ligands

6 Metal-ligand -bonding orbs,3 Metal-ligand -bonding orbs.

So optimum bonding at 18 e

Also large gap (oct) to metal-ligand * orbital, so rarely occupied

Exceptions:i) when ligands are not strong -acceptors andii) at extremes of TM rows

4p

4s

5 x 3d

Ligand -orbs

E

Ligand -orbs

-donor -donor/-acceptor

Metal carbonyls(compounds containing the MCO fragment)

L2

Coordination modes in TM carbonyls:• Always C-coordinated to TMs, and 2 e donor

O O O

C C C

M M M M MM

terminal (η1) doubly bridging (μ) triply bridging (μ3) e.g. [Cr(CO)6] e.g. [Co2(μ-CO)2(CO)6] e.g. [Co6(μ3-CO)4(CO)12]

(common form, and less commonalmost always in conjunctionwith M-M bonds)Also [Fe2(CO)9]

Preparation of TM carbonyls:

1. Direct combination:

– metal + CO (only works for Fe, Ni, and Ru)

2. Synthesis from a metal salt:

– metal salt + reducing agent + CO

Stable binary carbonyls (i.e. TM compounds with only CO ligands) are known for:

V Cr Mn Fe Co Ni

Mo Tc Ru Rh

W Re Os Ir

Group 6: Cr, Mo and W

• Synthesis by reduction of a metal salt:

– White, air stable volatile solid: octahedral:

• Cr-C 1.92 Å; IR: (CO) 2000 cm-1 (Point Group = Oh; t1u is the only IR active band expected)

• [Mo(CO)6] and [W(CO)6] have very similar properties

benzene

CrCl3 + Al + 6 CO [Cr(CO)6] + AlCl3

AlCl3

140oC, 300 atm

Group 8: Fe, Ru and Os• Iron: Synthesis is by direct combination:

200 atm

Fe + 5 CO [Fe(CO)5]

200oC

– [Fe(CO)5] is a yellow liquid, bp 103oC; structure: trigonal bipyramid;

– UV irradiation precipitates [Fe2(CO)9] (contains one Fe-Fe bond to keep 18 e)

• Ruthenium and osmium: Synthesis by reduction of a metal salt (e.g. via RuCl3 or OsO4)

– [Ru(CO)5] and [Os(CO)5] are unstable, since for these metals:

[M(CO)5] [M3(CO)12] + 3CO

– The M(CO)4 units in the final product are 16 e species, so 2 M-M bonds are needed to

allow each metal atom to achieve an 18 electron count - triangular

Group 10: Ni, Pd and Pt

• Nickel: Synthesis is by direct combination:

1 atm

Ni + 4 CO [Ni(CO)4]

25oC

(this is the basis of the Mond process for purification of Ni)

– Colourless liquid: bp 43 oC - very toxic (toxicity similar to HCN)

– Tetrahedral (Td) complex - Ni(0) is d10 - 18 e count; (CO) 2058 cm-1 (t2 expected from Group Theory)

• There are no binary carbonyls of Pd and Pt. Nearest equivalents are the [M(PR3)4] complexes or cluster anions, e.g. [Pt6(CO)12]2-

Group 9: Co, Rh and Ir

• Synthesis is by reduction of a metal salt:

130oC

2CoCO3 + 2H2 + 8CO [Co(CO)4]2 + 2CO2 + 2H2O

300 atm

– Co(CO)4 is a 17 e species, so it dimerises with a Co-Co bond (2 centre, 2 electron) to achieve 18 e count.

• There are 2 bridging CO groups - for e-counting purposes each bridging CO donates 1 electron to each metal.

– Heating causes loss of CO, forming [Co(CO)3]4 - 15 electron metal centres, requiring 3 Co-Co bonds - tetrahedron.

• M4(CO)12 is the smallest binary carbonyl for Rh and Ir – also made by reduction of a metal salt.

Group 7: Mn, Tc and Re

• Synthesis by reduction of a metal salt

CO (200 atm)

MnCl2 + Na/Ph2CO [Mn(CO)5]2

200oC, THF

(Ph2CO = dibenzophenone - forms a carbanion with sodium which reduces the metal)

– Electron count: Each Mn(CO)5 - with Mn(O) being d7 has 17 electrons, forming a 2-centre 2-electron Mn-Mn bond provides 18e count.

– (c.f. 2 CH3• H3C-CH3). d(Mn-Mn) = 2.92 Å. Staggered Mn(CO)5 units

• [Tc2(CO)10] and [Re2(CO)10] are also known and stable.

Group 5: V, Nb and Ta

• Synthesis is by reduction of a metal salt:

diglyme/ether

VCl3 + 3Na + 6CO [Na(diglyme)2]+[V(CO)6]-

120oC,200atm

(diglyme = MeOCH2CH2OCH2CH2OMe)

• [V-I(CO)6]- = a carbonylate anion - d6, oct. 18 electrons. (CO) 1859 cm-1

• Oxidation with HCl forms [V(CO)6] = 17 electrons, pyrophoric; μeff = 1.71 BM (t2g

5) – i.e. low spin (since CO is a strong ligand field)

• Nb2(CO)12 and Ta2(CO)12 have never been authenticated, but [M(CO)6]- are well known for Nb and Ta

Group 11: Cu, Ag and Au

• There are no stable binary carbonyls with Cu, Ag or Au

• Some complexes of M(I) (d10) with CO are known:

H2O

Cu + Cu2+ + CO [OC-Cu-OH2]+

• For Ag and Au, linear complexes known:

[OC-Ag-CO]+ (CO) = 2198 cm-1 (IR)

[OC-Au-Cl] (CO) = 2152 cm-1 (IR)

[OC-Au-CO]+ (CO) = 2217 cm-1 (IR)

• Virtually no -back bonding due to low energy of the d orbitals (and consequently high frequency IR bands, since the CO(*) orbital ~empty)

Metal carbonylates: (anionic metal carbonyls)

• Synthesis is by reduction of M-M bonds by Na (or as amalgam in Hg)

THF

Mn2(CO)10 + Na Na[Mn-I(CO)5]

• Further reduction by stronger reductant leads to:

THF

Mn2(CO)10 + (10x)Na Na3[Mn-III(CO)4]

– Mn(-III), d10. (CO) 1805w and 1670vs cm-1

• Base attack on a M-CO group:

Fe(CO)5 + 4 OH- [Fe-II(CO)4]2- + CO32- + 2 H2O

– Fe(-II), d10. (CO) 1771w and 1729vs cm-1

Metal hydrides (M-H):

1. Synthesis is by reduction of M-M bonds using H2:

[Co2(CO)10] + H2 [CoIH(CO)4] (pKa= 8.3)

2. Protonation of a carbonylate anion (with H2O or acid):

[Fe(CO)4]2- + H2O [Fe0H(CO)4]- + OH-

[FeoH(CO)4]- + H+ [FeIIH2(CO)4] (pKa = 11.4)

3. Base attack on a M-CO group:

Fe(CO)5 + 2 OH- [Fe0H(CO)4]- + HCO3-

Spectroscopic characterisation

IR spectroscopy - Functional groups & typical ranges:

Neutral carbonyls: (CO) for bridging groups 1650 - 1860 cm-1

(CO) for terminal groups 2125 - 1850 cm-1

(strong)

Nitrosyls: Linear M-NO+: (N-O) bands 1700 - 1900cm-1

Dinitrogen complexes: (N-N) bands 1900 - 2200cm-1 (medium

intensity)

Metal hydrides: (M-H) bands 2100 - 1800 cm-1 (weak)

Spectroscopic characterisationNMR spectroscopy - Functional groups:

Transition metal hydrides: characteristic 1H shift range: -5 to -30

e.g. (CO)5MnH -7.5

(CO)4FeH2 -11.1

(CO)4OsH2 -8.7

13C Shifts (vs. TMS as = 0):

Carbonyl ligands: Terminal: = 180 - 220

Bridging: = 220 - 280

Electron rich metals shift to higher

Comparison of -acceptor ligands

L3

Bonding in diatomics:

Electronegativity χP: C < N < O

N2 CO CN- NO+ ; RNC is pseudo-isoelectronic (RN like O); NO has 1 extra electron in *

2p

2s

2p

2s

N: 2s2 2p3 N: 2s2 2p3

Dinitrogen: N2

Carbon Monoxide: CO

C: 2s2 2p2 O: 2s2 2p4

2s

2p

2p

2s

*

*

* *

*

*Energy Energy

Dinitrogen frontier orbitals:

N N

N2 has 2 x degenerate Lowest

Unoccupied MOs (LUMOs) = * (mutually orthogonal)

Highest Occupied MO

(HOMO) = σ

N N

Carbon monoxide frontier orbitals:

C O

C O

2 x degenerate LUMOs = *(orthogonal)

Localised on C

HOMO = σ

Localised on C

MeNC 12

N2

8 CO

4 * eV 0

-4NO

-8

-12

-16

Comparison of -acceptor (-acid) ligands

•Interactions are strongest when orbital is localised on the donor atom.

•Donation is stronger from the higher energy occupied orbitals

•Acceptance is easier into lower lying unoccupied orbitals

•Cone angles of diatomics 90-95o

-acceptance: MeNC < N2 < CO << NO

σ-donation: MeNC > NO » CO > N2

Metal-CO bonding overview:

-donation: 2 electrons from HOMO – which is located mostly on C (hence M-C bond)

-backbonding: 2 * orbitals per CO – donation is from filled metal d orbitals

• Synergic bonding: i.e. the two forms of bonding are mutually enhancing

-acceptor ligands act as electron sinks - varying the extent of back donation

• In [M(CO)6], there are 9 M-L bonding orbitals occupied for 6 bonds bond order of 1.5

• In [M(CO)3L3] (where L is not a -acid ligand) then 3 M-L -bonds are localised on 3 M-CO bonds giving a M-CO bond order of 2.

Evidence for the synergic model (1):• Predicts that with increasing -acceptance:

– The M-C bond should shorten.

– The C-O bond should lengthen.

– The bond strength of the M-C bond will increase and hence the force constant of that bond might increase.

– The bond strength of the C-O bond will decrease and hence the force constant of that bond might decrease.

Evidence for the synergic model (2):• Note the following experimental constraints:

– the variation in C-O bond length is small compared to experimental error so the M-C bond length provides better evidence (X-ray crystallography)

– The (M-C ) modes mix with (M-C-O) and so the observed frequency is a reflection not only of the force constant KCO.

– The frequency range of (M-C ) is ~400 cm-1 - below the NaCl cut-off, so observable only with KBr or CsI windows in IR

– On the positive side - the (C-O) modes are very intense and occur in a region (~ 2000 cm-1) relatively free of other vibrations (except (M-H) )

• Hence the (CO) IR bands provide the easiest means of monitoring the relative degree of -backbonding into the * orbitals of CO

• The (NN) and (NO) bands will behave similarly in M-N2 and M-NO complexes.

(CO) energies in metal carbonyls:• In free CO, (CO) = 2143 cm-1

• For a diatomic molecule: (Hz) = (1/2)√(K/m)

– where m (reduced mass) = m1m2/(m1+m2) and K = the force constant of the bond

• Consider a series of isoelectronic complexes – we see an increase in nuclear charge from V(-I) to Mn(+I), or from Fe(-II) to Ni(0) and hence increasing Zeff of the 3d orbitals. So less available for -back bonding:

Experimentally:

[Mn(CO)6]+ [Cr(CO)6] [V(CO)6]-

(CO) 2096 2000 1859(MC) 416 449 460

[Ni(CO)4] [Co(CO4]- [Fe(CO)4]2-

(CO) 2058 1883 1729

Increasing back bonding to CO from left to right

CO bond order and (CO) from left to rightMC bond order and (MC) from left to right

Effect of CO substitution:

• With the exception of PF3, phosphines are stronger donors and weaker acceptors than CO. So substitution of a CO by a phosphine should increase the electron density on a metal and thus increase the -back bonding to the remaining CO groups (compare the A1 modes across the series).

– CO vibrational modes for [Ni(CO)4-nLn], L = P(OMe)3

• The (CO) frequencies are good guides to the mean donor/acceptor ability of the other ligands (L = PR3) in the complex.

– E.g. [Ni(CO)3L], A1 modes:

[Ni(CO)]4 [Ni(CO)3L] [Ni(CO)2L2][Ni(CO)L3]

Point Td C3v C2v C3v

group:

(CO) A1 2131 A1 2082 A1 2025 A1 1964 T2 2058 E 2006 B1 1965

(MC) A1 368 A1 420 A1 458 A1 499 T2 423 E 455 B1 495

R (CO) (A1) -properties -propertiesof L (PR3) of L (PR3)

C(CH3)3 2056 -donation -acceptanceCH3 2064 decreasing increasingC6H5 2069OCH3 2079OC6H5 2085F 2111

As (CO) increases, this indicates that PR3 is pushing less e-density

onto Ni; i.e. PR3 is a weaker -donor or a better -acceptor

Metal nitrosyls:• Sources of NO in a complex are:

– NO+BF4-, NOCl, NO2 - all with N as NIII; and NO with N as NII

NO+ + e = NO (Eo = 1.28 V vs SCE in MeCN)

(NO): ~2250 cm-1 1840 cm-1

• NO can reduce the metal formally to form NO+ on coordination and displace CO:

[Fe(CO)5] + 2 NO [Fe-II(CO)2(NO)2] + 3 CO

d8, 18 e d10, 18e (C2v)

(CO) 2087, 2038 cm-1;

(NO) 1817, 1780 cm-1

[Cr(CO)6] + 4 NO [Cr-IV(NO)4]

d6, 18 e d10, 18e (Td)

(NO) 1726 cm-1

• Important reaction type inter-converting nitrosyls and nitrites:

[(CN)5Fe(NO)]2- + OH- [(CN)5Fe{NO(OH)}]3- [(CN)5Fe(NO2)]4-

Metal nitrosyls: alternative structure• In some complexes the M-N-O angle is highly bent:

[IrIIICl(NO-)(CO)(PPh3)2]+:

5d6 2e 2e 2e 4e = 16 electrons

• the NO is now considered as a [NO]- with the extra 2e in a lone pair on an sp2 N atom.

– Metal forms a σ-bond by donation from the N, and there is -back-bonding in one plane:

Ir ClP

OCP

NO

-bonding interaction -bonding interaction

M

N O

M

N O

NB – within CHEM2009 you should treat TM-NO complexes as containing

NO+ (2e donor) unless specifically told otherwise

For background information:

Metal dinitrogen complexes:• Direct synthesis from N2:

[RuII(NH3)5(OH2)]2+ + N2 H2O + [Ru(NH3)5N2]2+

d6, 18e d6, 18e

(NN): 2331 cm-1 (Raman) 2150 cm-1 (IR)

• Reduction in the presence of N2:

[FeCl2{Et2PCH2CH2PEt2}2] + N2 + Na(naphthalenide) [Fe(N2){Et2PCH2CH2PEt2}2]

• Reaction with acids (Functional model of nitrogenase):

[WCl4(PMe2Ph)3] + N2 + Mg + PMe2Ph [W(N2)2(PMe2Ph)4]

H+

NH3

Metal-carbon σ-bonds

L4

Synthesis of metal alkyls (1):• From substitution of anionic ligands:

– the M-C bond involves alkyl acting as a 2 e donor:

– metal halide + carbanion source (M’ = Li, Mg, Al, Zn, Hg)

ZrCl4 + 4 PhCH2MgCl [Zr(CH2Ph)4]

TiCl4 + Al2Me6 [MeTiCl3]

• Metal hydride + alkene: homologation of an anionic ligand (insertion)

trans-[(Et3P)2PtHCl] + C2H4 trans-[(Et3P)2Pt(C2H5)Cl]

NB there is no change in oxidation state during this reaction

Synthesis of metal alkyls (2):• From alkyl electrophiles:

– Carbonylate anion + alkyl halide

Na+[Mn-I(CO)5]- + CH3I [CH3MnI(CO)5] + NaI

• From 16e complex + alkyl halide:

[IrI(CO)L2Br] + CH3Cl [IrIII(CO)L2Br(CH3)Cl]

Increase in oxidation state of II – formally reducing a carbenium ion (carbocation) to a carbanion (oxidative addition- see later)

Stability of metal alkyls (1):

• Thermodynamic stability:

– Comparable to main group alkyls

– Decreases with the number of alkyl groups

– Bond strength increases down the group (unlike main group alkyls)

(d orbitals become less contracted as the TM Group is descended, hence giving better overlap with ligand orbitals)

Main Group Alkyl D(kJ/mol) TM Alkyl D (kJ/mol)

CMe4 358 Ti(CH2tBu)4 198

SiMe4 311 Zr(CH2tBu)4 249

GeMe4 249 Hf(CH2tBu)4 266

SnMe4 217 TaMe5 261

PbMe4 152 WMe6 160

Stability of metal alkyls (2):

• Kinetic stability:

– Often much lower than main group alkyls due to different decomposition mechanisms:

• Main group decomposition (homolytic fission – radicals) (PbEt4 is quite stable)

PbEt4 (100o C) PbEt3 + Et

• TM alkyl decomposition (β-hydride elimination) (TiEt4 is unknown)

R3Ti-CH2CH3 R3TiH + CH2=CH2 (see later)

- Increase stability when no β-H present (e.g. R = CH3, CH2CMe3, CH2Ph, CH2SiMe3)

- Increase stability by using more bulky ligands

• More electron rich alkyls as donors and block dinuclear decomposition mechanisms

i.e. TiMe4 is stable to –40oC but the much bulkier Ti(CH2CMe3)4 is stable > 100oC

NMR SpectroscopicCharacterisation

Spectroscopic characterisationof M-alkyls/alkenes/arenes:

NMR spectroscopy - Functional groups:

1H NMR

• Metal alkyls: Shifts are a function of electronegativity (oxidation state):

• electron rich metals shift to lower : Range 5 to -1 ppm

• Arenes and alkenes:

• Upfield (lower ) shifts on coordination ~ 0.5 to 4 ppm;

• electron rich metals shift to lower

13C NMR

• Metal alkyls: shifts function of electronegativity (oxidation state):

• electron rich metals shift to lower : Range 50 to -20

• Arenes and alkenes: Upfield (lower ) shifts on coordination of ~ 5 to 80 ppm

• electron rich metals shift to lower :

13C{1H} NMR Example: [Re{C(O)CH3}(CO)5]

ReCOOC

COOCCO

O CH3

/ppm

CH3CO trans

4 x CO

C (acyl)

245 55200 150 100

Metal-carbon -bonds

Metal alkene complexes:• Synthesis by substitution reactions of 2 e donor ligands:

K2[PtCl4] + C2H4 K+[(-C2H4)PtIICl3]-.H2O + KCl

Re(CO)5Cl + C2H4 + AlCl3 [ReI(CO)5(-C2H4)]+AlCl4-

• Centre point of the double bond acts as the ligand binding point

• Alkene is face-on to the metal

– An extreme view would be to consider it as a

[C2H4]2- bidentate metal-di-alkyl, but generally

the actual bonding is more subtle than that

description.

Bonding of ethene to a transition metal (1):

• Highest Occupied Molecular Orbital (HOMO) can act as a 2-electron donor.

• Lowest Unoccupied Molecular Orbital (LUMO) can act as a -acceptor.

Bonding of ethene to a transition metal (2)• The orbital on ethene acts as a σ-donor to the

metal;

• the * orbital on ethene can be a acceptor from the metal.

C

C

M

-donation(2e donor)

Back-bonding to empty orbital

C

CM

Bonding of ethyne (RCCR) to a TM:

• One orbital on ethyne acts as a σ-donor to the metal and one * orbital on ethene can be a -acceptor from the metal.

HOMO as 2 e donor HOMO with possible weak donation

LUMO as acceptor LUMO as acceptor

with weak -interaction

• Generally alkynes are 2 electron donors (– occasionally 4 e donors to early transition metals with low electron density)

Alkyne bonding orbital interactions

-donation from filled alkyne -orbital

-acceptance to vacant *

alkyne orbital

-donation from filled alkyne orbital

-type back bonding to vacant *

alkyne orbital

TM-alkyne bonding• The bonding of an alkyne to a TM complex is similar to that of

an alkene complex.

• Alkynes tend to be more electropositive and therefore tend to bind more tightly to a TM than alkenes.

• In fact, alkynes will often displace alkenes.

• The primary difference in bonding between alkenes and alkynes is that an alkyne can act as either a 2 or 4 electron donor.

• Alkynes have two sets of mutually orthogonal bonds. We can bind one of these to the TM in a -donor fashion (A) as we did for alkenes, including a -backbond (B).

• The orthogonal set can also bind in a -donor fashion using an orthogonal metal d-orbital (C).

• The back-donation to the anti-bonding orbital (D) is a -bond, and the degree of overlap is small (the contribution of D to the bonding of alkynes is minimal at best).

Metal-Carbon bonds – arenes and cyclopentadienyl ligands L5

Arenes as ligands (1):• The 6 electrons in benzene can also act as a ligand –

‘as if a delocalised tri-alkene ligand’

nBu2O, THF

[Cr(CO)6] + C6H6 [Cr(6-C6H6)(CO)3] + 3 CO

Arene is face-on to metal so 6 -electrons act as donors (molecule has C3v symmetry).

Can consider as an octahedral complex with the benzene a facially coordinating ligand (effectively occupying 3 coordination sites).

COOC CO

Cr

Arenes as ligands (2):• With 2 benzene ligands – can form a sandwich complex

– [M(C6H6)2] structures are mostly stable as 18 electron species

Both arenes are face-on to metal so6 -electrons act as donors.

Rings are eclipsed in the crystalstructure (D6h symmetry).

Can consider as an octahedralcomplex with the benzenes as facially coordinating ligand

E. O. Fischer, 1955 - rational synthesis and structure; Nobel Prize 1973

6 electrons

3d6

6 electrons

= 18 electrons

Cr 0

Charged arenes as ligands (1):• The 6 electrons in the cyclopentadienyl anion (C5H5)- can

also act as a strong ligand

5-C5H5- ring face-

on to metal: aromatic anion donates 6 -electrons to metal

[Mo(CO)6]

MoCOOC CO

(+ Na)

_

+ H+

_

+ 3 CO

cyclopentadiene, C5H6

cyclopentadienyl, C5H5-

Point group = C3v; CO = a1 + e -both IR active (2 CO

stretches in IR)

Charged arenes as ligands (2):• Two cyclopentadienyl anions (C5H5)- (often abbreviated as

Cp-) can also form sandwich complexes: metallocenes

– these exist from M = V to Ni (15 to 20 electrons) due to the stability of the ion triplet.

Synthesis: P L Pauson, 1951.Proposal of sandwich structure: G Wilkinson, 1951: Nobel Prize

1973

Ferrocene: [Fe(5-C5H5)2]

Cyclopentadienyl rings both face-on to the metal

Nearly staggered in crystal

Fe

6 e

d6

6 e

= 18 e

II

Case studies/examples:

• Determine the NVE associated with the following species and sketch their structures:

– [FeH2(CO)4] 1H NMR: = -8.4

– [Ru3(CO)12]

– [Pt(tBu)Cl(PMe3)2]

– [(5-C5H5)Fe(CO)2I] (CO) = 2055, 1982 cm-1

For the reaction sequence below, use the data provided to identify compounds A-D and sketch their structures.

[MnCl(CO)5] + PMe3 A + B

C + D

CH2Cl2, reflux

(i) AgBF4 / Me2CO; (ii) PMe3

A: %C = 33.1; %P = 19.0; %Mn = 16.8; IR: 2028, 1959, 1917 cm-1; (31P) = -17 (s).

B: m/z = 28

C: %C = 31.8; %P = 20.5; %Mn = 12.1; IR: 2020, 1950 cm-1; (31P) = -10 (s).

D: white solid; %Cl = 24.8

[PMe3: (31P) = -62]