1. Chapter 13 Lecture 1 Organometallic Ligands and Bonding
Organometallic Basics
An Organometallic Complex contains at least one M—C bond
Includes ligands: CO, NO, N2, PR3, H2
Doesn’t include CN- (classical coordination chemistry ligand)
Both s and p bonding between M and C occur
History
Zeise’s Salt synthesized in 1827 = K[Pt(C2H4)Cl3] • H2O
Confirmed to have H2C=CH2 as a ligand in 1868
Structure not fully known until 1975
Ni(CO)4 synthesized in 1890
Grignard Reagents (XMgR) synthesized about 1900
Accidentally produced while trying to make other compounds
Utility to Organic Synthesis recognized early on

2. 4) Ferrocene synthesized in 1951
a) Modern Organometallic Chemistry begins with this discovery
b) Many new ligands and reactions produced ever since
5) Organometallic Chemistry has really been around for millions of years
a) Naturally occurring Cobalimins contain Co—C bonds
b) Vitamin B12

3. Ligands and Nomenclature
Common Organic Ligands
Binding Modes
Bridging is possible with organometallic ligands

4. Different numbers of atoms of the organometallic ligand may be involved in bond and is called the “Hapticity” of the ligand
The 18-electron Rule
Counting Electrons
The octet rule governs organic and simple ionic compounds: s + 3p orbital
The 18-electron rule governs organometallics (with many exceptions)
s + 3p + 3d orbitals
Donor ligands provide the electrons other than the d-electrons

5. 3) The “Donor Pair” method of electron counting (Method A in your book)
a) Common organometallic ligands are assigned an electron count and charge
b) The charge on ligands helps determine d-electron count of metal
c) Add up all electrons from Metal d orbitals and ligands to find total e- count
(isonitrile or isocyanide)
(oxo, sulfido)
(nitrido)

6. Examples of Electron Counting
Cr(CO)6
Total charge on ligands = 0, so charge on Cr = 0, so Cr = d6
6 CO ligands x 2 electrons each = 12 electrons
Total of 18 electrons
(h5-C5H5)Fe(CO)2Cl
Total charge on ligands = 2-, so Fe2+ = d6
(h5-C5H5- = 6) + (2CO x 2 = 4) + (Cl- = 2) = 12 electrons
Charged complex: [Mn(CO)6]+
Total ligand charge = 0, so Mn+ = d6
12 electrons from 6 CO ligands gives a total of 18 electrons
M—M Bond: (CO)5Mn—Mn(CO)5
Each bond between metals counts 1 electron per metal: Mn—Mn = 1 e-
Total ligand charge = 0, so Mn0 = d7
5 CO ligands per metal = 10 electrons for a total of 18 electrons per Mn

7. Justification for and exceptions to the 18-electron Rule
MO Theory predicts that 18 electrons fill bonding orbitals
This number is more stable than more (filling antibonding orbitals) or less Do

8. When is the 18-electron rule most valid?
a) With octahedral complexes of large Do.
Ligands are good s-donors and good p-acceptors (CO)
Exceptions to the 18-electron rule are common
Weak field ligands with small Do make filling eg* possible ( > 18e-)
p-donor ligands can make t2g antibonding ( < 18 e-)

9. 5). Square Planar Complexes (d8) follow a 16-electron rule
5) Square Planar Complexes (d8) follow a 16-electron rule. 18 electrons would destabilized the complexes by filling the high energy dx2-y2 orbital.

10. Carbonyl Complexes (CO)
Bonding
Review of CO Molecular Orbitals
HOMO resides mostly on C = s-donation
LUMO resides mostly on C = p-acceptance
Reinforce each other and provide strong bonding
Bonding of CO to a Metal

11. Characteristics of CO complexes
Infrared Spectroscopy
Free CO stretch n = 2143 cm-1
Cr(CO)6 CO stretch n = 2000 cm-1 because p-back donation from metal weakens the CO bond by adding e- to antibonding p* orbital
Negative charge on complex further weakens CO bond:
[V(CO)6]- n = 1858 cm [Mn(CO)6]+ n = 2095 cm-1
d) Bridging CO further weakened by extra p-back donation (e- count = 1/M)
X-Ray Crystallography
Free CO bond length = pm
M—CO carbonyl bond length = 115 pm

12. Synthesis and Reactions of CO Complexes
Carbonyl complexes of most metals exist.
Most obey the 18-electron rule
Bridging decreases down the periodic table as d-orbitals become larger.
Synthesis
Direct reaction: Ni CO Ni(CO)4 Toxic, used to purify Ni
Reductive Carbonylation: CrCl3 + CO + Al Cr(CO)6 + AlCl3
Thermal/Photochemical: 2 Fe(CO)5 + hn Fe2(CO)9 + CO
Reactions: useful for the synthesis of other compounds by substitution of CO
Cr(CO) PPh Cr(CO)5(PPh3) CO
Re(CO)5Br en Re(CO)3(en)Br CO

14. Ligands Similar to CO
CN- (cyanide) is isolectric to CO
Stronger s-donor and slightly weaker p-acceptor than CO
More stable with M+ due to –1 charge than M0 (which favors CO)
Considered a classical ligand rather than organometallic for this reason
NN (dinitrogen) is isoelectric to CO
Weaker s-donor and weaker p-acceptor than CO, so doesn’t bind well
Very important in Nitrogen Fixation, so much research centers on complexes
NO+ (nitrosyl) is isoelectric to CO
Similar to CO in s-donor and p-acceptor properties
Electron counting scheme considers linear NO complexes as 2 e- NO+
Electron counting scheme considers bent NO complexes as 2 e- NO-

15. Hydride and Dihydrogen Complexes
Hydride Complexes
M—H bonding is s-donation only from H- (2 electron, -1 charge)
Synthesis
Reaction with H2: Co2(CO) H HCo(CO)4
Reduction of carbonyl complex followed by addition of H+
Co2(CO) Na Na[Co(CO)4]
2 Na[Co(CO)4] + H HCo(CO)4
Dihydrogen Complexes
First Complex characterized in Mo(CO)3(PR3)(H2)
Bonding
s-donation from H2 s molecular orbital
p-acceptance from H2 s* molecular orbital
H—H bond is weaker and longer than free H2 (82-90 pm vs. 74 pm)
Electron-rich metals can completely rupture the H2 bond by p-back donation
Other good p-acceptor ligands on the metal helps stabilize the H2—M complex