Transcript Chapter 1 Structure and Bonding

Chapter 13 Lecture 1 Organometallic Ligands and Bonding
I.
Organometallic Basics
A.
An Organometallic Complex contains at least one M—C bond
1) Includes ligands: CO, NO, N2, PR3, H2
2) Doesn’t include CN- (classical coordination chemistry ligand)
3) Both s and p bonding between M and C occur
B.
History
1) Zeise’s Salt synthesized in 1827 = K[Pt(C2H4)Cl3] • H2O
a) Confirmed to have H2C=CH2 as a ligand in 1868
b) Structure not fully known until 1975
2)
Ni(CO)4 synthesized in 1890
3)
Grignard Reagents (XMgR) synthesized about 1900
a) Accidentally produced while trying to make other compounds
b) Utility to Organic Synthesis recognized early on
4)
5)
Ferrocene synthesized in 1951
a)
Modern Organometallic
Chemistry begins with this
discovery
b)
Many new ligands and reactions
produced ever since
Organometallic Chemistry has really
been around for millions of years
a)
Naturally occurring Cobalimins
contain Co—C bonds
b)
Vitamin B12
II.
Ligands and Nomenclature
A.
Common Organic Ligands
B.
Binding Modes
1) Bridging is possible with organometallic ligands
2)
Different numbers of atoms of the organometallic ligand may be involved in bond
and is called the “Hapticity” of the ligand
III. The 18-electron Rule
A.
Counting Electrons
1) The octet rule governs organic and simple ionic compounds: s + 3p orbital
2)
The 18-electron rule governs organometallics (with many exceptions)
a) s + 3p + 3d orbitals
b) Donor ligands provide the electrons other than the d-electrons
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)
B.
Examples of Electron Counting
1) Cr(CO)6
a) Total charge on ligands = 0, so charge on Cr = 0, so Cr = d6
b) 6 CO ligands x 2 electrons each = 12 electrons
c) Total of 18 electrons
2)
(h5-C5H5)Fe(CO)2Cl
a) Total charge on ligands = 2-, so Fe2+ = d6
b) (h5-C5H5- = 6) + (2CO x 2 = 4) + (Cl- = 2) = 12 electrons
c) Total of 18 electrons
3)
Charged complex: [Mn(CO)6]+
a) Total ligand charge = 0, so Mn+ = d6
b) 12 electrons from 6 CO ligands gives a total of 18 electrons
4)
M—M Bond: (CO)5Mn—Mn(CO)5
a) Each bond between metals counts 1 electron per metal: Mn—Mn = 1 eb) Total ligand charge = 0, so Mn0 = d7
c) 5 CO ligands per metal = 10 electrons for a total of 18 electrons per Mn
C.
Justification for and exceptions to the 18-electron Rule
1) MO Theory predicts that 18 electrons fill bonding orbitals
2) This number is more stable than more (filling antibonding orbitals) or less
Do
3)
When is the 18-electron rule most valid?
a) With octahedral complexes of large Do.
b) Ligands are good s-donors and good p-acceptors (CO)
4)
Exceptions to the 18-electron rule are common
a) Weak field ligands with small Do make filling eg* possible ( > 18e-)
b) p-donor ligands can make t2g antibonding ( < 18 e-)
5)
Square Planar Complexes (d8) follow a 16-electron rule. 18 electrons would
destabilized the complexes by filling the high energy dx2-y2 orbital.
IV. Carbonyl Complexes (CO)
A.
Bonding
1)
Review of CO Molecular Orbitals
a) HOMO resides mostly on C = s-donation
b) LUMO resides mostly on C = p-acceptance
c) Reinforce each other and provide strong bonding
2)
Bonding of CO to a Metal
B.
Characteristics of CO complexes
1) Infrared Spectroscopy
a) Free CO stretch n = 2143 cm-1
b) 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
c) Negative charge on complex further weakens CO bond:
[V(CO)6]- n = 1858 cm-1 [Mn(CO)6]+ n = 2095 cm-1
d) Bridging CO further weakened by extra p-back donation (e- count = 1/M)
2)
X-Ray Crystallography
a) Free CO bond length = 112.8 pm
b) M—CO carbonyl bond length = 115 pm
C.
Synthesis and Reactions of CO Complexes
1) Carbonyl complexes of most metals exist.
a) Most obey the 18-electron rule
b) Bridging decreases down the periodic table as d-orbitals become larger.
2)
3)
Synthesis
a) Direct reaction: Ni + 4 CO
Ni(CO)4
b)
Reductive Carbonylation: CrCl3 + CO + Al
c)
Thermal/Photochemical: 2 Fe(CO)5 + hn
Toxic, used to purify Ni
Cr(CO)6 + AlCl3
Fe2(CO)9 + CO
Reactions: useful for the synthesis of other compounds by substitution of CO
a)
Cr(CO)6 + PPh3
Cr(CO)5(PPh3) + CO
b)
Re(CO)5Br + en
Re(CO)3(en)Br + 2 CO
D.
Ligands Similar to CO
1) CN- (cyanide) is isolectric to CO
a) Stronger s-donor and slightly weaker p-acceptor than CO
b) More stable with M+ due to –1 charge than M0 (which favors CO)
c) Considered a classical ligand rather than organometallic for this reason
2)
NN (dinitrogen) is isoelectric to CO
a) Weaker s-donor and weaker p-acceptor than CO, so doesn’t bind well
b) Very important in Nitrogen Fixation, so much research centers on complexes
3)
NO+ (nitrosyl) is isoelectric to CO
a) Similar to CO in s-donor and p-acceptor properties
b) Electron counting scheme considers linear NO complexes as 2 e- NO+
c) Electron counting scheme considers bent NO complexes as 2 e- NO-
V.
Hydride and Dihydrogen Complexes
A.
Hydride Complexes
1) M—H bonding is s-donation only from H- (2 electron, -1 charge)
2) Synthesis
a) Reaction with H2: Co2(CO)8 + H2
2 HCo(CO)4
b)
B.
Reduction of carbonyl complex followed by addition of H+
Co2(CO)8 + 2 Na
2 Na[Co(CO)4]
2 Na[Co(CO)4] + H+
HCo(CO)4
Dihydrogen Complexes
1) First Complex characterized in 1984 Mo(CO)3(PR3)(H2)
2) Bonding
a) s-donation from H2 s molecular orbital
b) p-acceptance from H2 s* molecular orbital
c) H—H bond is weaker and longer than free H2 (82-90 pm vs. 74 pm)
3)
Electron-rich metals can completely rupture the H2 bond by p-back donation
4)
Other good p-acceptor ligands on the metal helps stabilize the H2—M complex