Download Types of bonds possible from Ligands

Types of bonds possible from Ligands
Language: All bonds are coordination or coordinative
Remember that all of these bonds are weaker than normal organic
bonds (they are dative bonds)
Simple ligands e.g. CH3-, Cl-, H2 give σ bonds
π systems are different e.g. CO is a σ donor and π acceptor
Bridging ligands can occur two metals
Metal-metal bonds occur and are called δ bonds – they are weak
and are a result of d-d orbital overlap
Alkyl ligands:
Transition metal alkyl complexes important for catalysts e.g.
olefin polymerization and hydroformylation thermodynamic
Problem is their weak kinetic stability
(Thermally fine: M-C bond dissociation energies are typically 4060 kcal/mol with 20-70 kcal/mol)
Simple alkyls are sigma donors, that can be considered to donate one
or two electrons to the metal center depending on which electron
counting formalism you use
Synthesis of Metal Alkyl Complexes
1. Metathetical exchange using a carbon nucleophile (R-).
Common reagents are RLi, RMgX (or R2Mg), ZnR2, AlR3, BR3,
and PbR4. Much of this alkylation chemistry can be understood
with Pearson's "hard-soft" principles
2. Metal-centered nucleophiles (i.e. using R+ as a reagent) Typical
examples are a metal anion and alkyl halide (or pseudohalogen).
for example:
NaFp + RX
Fp-R + NaX
[Fp = Cp(CO)2Fe]
3. Oxidative Addition. This requires a covalently
unsaturated, low-valent complex (16 e- or less). A classic
example:
4. Insertion- To form an alkyl, this usually involves an olefin
insertion. The simplest generic example is the insertion of
ethylene into an M-X bond, i.e.
M-X + CH2CH2
M-CH2CH2-X
Carbonyl Complexes
Bonding of CO
Electron donation of the lone
pair on carbon σ This electron
donation makes the metal
more electron rich compensate for this increased
electron density, a filled metal
d-orbital may interact with the
empty π* orbital on the
carbonyl ligand
π-backbonding or πbackdonation or synergistic
bonding
Similar for alkenes,
acetylenes, phosphines, and
dihydrogen.
What stabilizes CO complexes is M→C π–bonding
The lower the formal charge on the metal ion the more willing it
is to donate electrons to the π–orbitals of the CO
Thus, metal ions with higher formal charges, e.g. Fe(II) form CO
complexes with much greater difficulty than do zero-valent
metal ions
For example Cr(O) and Ni(O), or negatively charged metal ions
such as V(-I)
In general to get a feeling for stability examine the charges on
the metals
Syntheses of metal carbonyls
Metal carbonyls can be made in a variety of ways.
For Ni and Fe, the homoleptic or binary metal carbonyls can be made by
the direct interaction with the metal (Equation 1).
In other cases, a reduction of a metal precursor in the presence of CO
(or using CO as the reductant) is used (Equations 2-3).
Carbon monoxide also reacts with various metal complexes, most
typically filling a vacant coordination site (Equation 4) or performing a
ligand substitution reactions (Equation 5)
Occasionally, CO ligands are derived from the reaction of a coordinated
ligand through a deinsertion reaction (Equation 6)
Synthesis of carbonyl complexes
Direct reaction of the metal
– Not practical for all metals due to need for
harsh
conditions (high P and T)
– Ni + 4CO ÆNi(CO)4
– Fe + 5CO ÆFe(CO)5
Reductive carbonylation
– Useful when very aggressive conditions
would be
required for direct reaction of metal and CO
» Wide variety of reducing agents can be used
– CrCl3+ Al + 6CO Æ AlCl3 + Cr(CO)6
– 3Ru(acac)3 + H2 + 12CO Æ Ru3(CO)12 +
Main characterization methods:
• Xray diffraction ⇒ (static) structure ⇒ bonding
• NMR ⇒ structure en dynamic behaviour
• EA ⇒ assessment of purity
• (calculations)
Useful on occasion:
• IR
• MS
• EPR
Not used much:
• GC
• LC
IR spectra and metal-carbon bonds
The υCO stretching frequency of the coordinated CO is
very informative
Recall that the stronger a bond gets, the higher its
stretching frequency
M=C=O (C=O is a double bond) canonical structure
Lower the υCO stretching frequency as compared to the MC≡O structure (triple bond)
Note: υCO for free CO is 2041 cm-1)
[Ti(CO)6]2- [V(CO)6]- [Cr(CO)6] [Mn(CO)6]+ [Fe(CO)6]2+
υCO
1748
1858
1984
increasing M=C double
bonding
2094
2204 cm-1
decreasing M=C double
bonding
Bridging versus terminal carbonyls
Bridging CO groups can be regarded as having a double
bond C=O group, as compared to a terminal C≡O, which is
more like a triple bond:
~ double bond
~ triple bond
M
M-C≡O
C=O
M
terminal carbonyl
(~ 1850-2125 cm-1)
the C=O group
in a bridging
carbonyl is more
like the C=O in
a ketone, which
typically has
υC=O = 1750 cm-1
bridging carbonyl
(~1700-1860 cm-1)
Bridging CO between 1700 and 2200 cm-1
Bridging versus terminal carbonyls in [Fe2(CO)9]
O
OC
OC
Fe
OC
CO
C
Fe
C
C O
CO
CO
O
terminal
carbonyls
bridging
carbonyls
Summary
1. As the CO bridges more metal centers its stretching
frequency drops – same for all π ligands
– More back donation
2. As the metal center becomes increasingly electron rich the
stretching frequency drops
Alkene ligands
Dewar-Chatt-Duncanson model
The greater the electron density
back-donated into the π* orbital
on the alkene, the greater the
reduction in the C=C bond order
Stability of alkene complexes
also depends on steric factors
as well
An empirical ordering of relative
stability would be:
tetrasubstituted < trisubstituted
< trans-disubstituted < cisdisubstituted < monosubstituted
< ethylene
Alkyne ligands:
Similar to alkenes
Alkynes tend to be more
electropositive-bind
more tightly to a
transition metal than
alkenes -alkynes will
often displace alkenes
Difference is 2 or 4
electron donor
sigma-type fashion (A)
as we did for alkenes,
including a pi-backbond
(B)
The orthogonal set can
also bind in a pi-type
fashion using an
orthogonal metal dorbital (C)
The back-donation to the antibonding orbital (D) is a deltabond-the degree of overlap is quite small - contribution of D to
the bonding of alkynes is minimal
The net effect π-donation - alkynes are usually nonlinear in TM complexes
Resonance depict the bonding of an alkyne.
I is the metallacyclopropene resonance form
Support for this versus a simple two electron donor,
II, can be inferred from the C-C bond distance as
well the R-C-C-R angles
III generally does not contribute to the bonding of
alkyne complexes.
Ally ligands:
Allyl ligands are ambidentate ligands that can bind in both a
monohapto and trihapto form The trihapto form can be
expressed as a number of difference resonance forms as
shown here for an unsubstituted allyl ligand: Important
applications
Dihydrogen Ligands:
Metal is more electropositive than hydrogen
Hydrogen acts as a two electron sigma donor to the metal center.
The complex is an arrested intermediate in the oxidative addition of
dihydrogen
How does this affect the oxidation state of the
meta?
Dihydrogen complexes Bonding is “simple” a 3C-2electron bond.
H2 - neutral two electron sigma donor
One could also describe a back-donation of electrons from a filled
metal orbital to the sigma-* orbital on the dihydrogen
Electronic Attributes of Phosphines
Like that of carbonyls
As electron-withdrawing sigma-donating capacity decreases
At the same time, the energy of the π-acceptor (sigma-*) on phosphorous
is lowered in energy, providing an increase in backbonding ability.
Therefore, range of each capabilities -tuning
Rough ordering -CO stretching frequency indicator
low CO stretching frequency- greater backbonding to M
Experiments such as this permit us to come up with the following
empirical ordering:
Cone Angle (Tolman)
Steric hindrance:
Phosphine
Ligand
Cone
Angle
A cone angle of 180 degrees effectively protects (or covers)
one half of the coordination
sphere of the metal complex
PH3
87o
PF3
104o
P(OMe)3
107o
PMe3
118o
PMe2Ph
122o
PEt3
132o
PPh3
145o
PCy3
170o
P(t-Bu)3
182o
P(mesityl)3
212o
You would expect a dissociation event
to occur first before any other reaction
-steric bulk (rate is first order
-increasing size)
This will also have an effect on
activity for catalysts
N.B. “flat” can slide past each other
For example Wilkinson's catalyst
(more later)
Has a profound effect on the
reactivity!
18 Electron Rule (Sidgwick, 1927)
•
•
OM chemistry gives rise to many “stable” complexes - how
can we tell by a simple method
Every element has a certain number of valence orbitals:
1 { 1s } for H
4 { ns, 3´np } for main group elements
9 { ns, 3´np, 5´(n-1)d } for transition metals
s
dxy
px
dxz
py
dyz
pz
dx2-y2
dz2
•
•
•
Therefore, every element wants to be surrounded
by 2/8/18 electrons
– For main-group metals (8-e), this leads to the standard Lewis
structure rules
– For transition metals, we get the 18-electron rule
Structures which have this preferred count are called
electron-precise
Every orbital wants to be “used", i.e. contribute to binding an
electron pair
The strength of the preference for electron-precise structures depends
on the position of the element in the periodic table
For early transition metals, 18-e is often unattainable for steric
reasons - the required number of ligands would not fit
• For later transition metals, 16-e is often quite stable (square-planar
d8 complexes)
• Addition of 2e- from 5th ligand converts complex to 5 CN 18e- ,
marginally more stable
•
Predicting reactivity
14 e
- C2H4
(C2H4)2PdCl2
16 e
CO
dissociative
(C2H4)PdCl2
CO
?
(C2H4)2(CO)PdCl2
associative
(C2H4)(CO)PdCl2
- C2H4
18 e
Most likely associative
16 e
Predicting reactivity
16 e
- CO
18 e
Cr(CO)6
MeCN
dissociative
Cr(CO)5
MeCN
?
Cr(CO)6(MeCN)
associative
Cr(CO)5(MeCN) 18 e
- CO
20 e (Sterics!)
Most likely dissociative
N.B. How do you know a fragment forms a covalent or a
dative bond?
•
•
•
Chemists are "sloppy" in writing structures. A "line" can
mean a covalent bond, a dative bond,
recognise/understand the bonding first
Use analogies ("PPh3 is similar to NH3").
Rewrite the structure properly before you start counting.
Cl
PPh3
Cl
Pd
covalent
bond
1e
dative
bond
"bond" to the
allyl fragment
PPh 3
2e
Pd
3e
Pd =
Cl⎯ =
P→ =
allyl =
10
1
2
3
+ ⎯⎯
e-count 16
"Covalent" count: (ionic method also useful)
1. Number of valence electrons of central atom.
• from periodic table
2. Correct for charge, if any
• but only if the charge belongs to that atom!
3. Count 1 e for every covalent bond to another atom.
4. Count 2 e for every dative bond from another atom.
• no electrons for dative bonds to another atom!
5. Delocalized carbon fragments: usually 1 e per C (hapticity)
6. Three- and four-center bonds need special treatment
7. Add everything
N.B. Covalent Model:
18 = (# metal electrons + # ligand electrons) - complex charge
The number of metal electrons equals it's row number (i.e., Ti = 4e, Cr
= 6 e, Ni = 10 e)
Hapto (η) Number (hapticity)
For some molecules the molecular formula provides insufficient
information with which to classify the metal carbon interactions
The hapto number (η) gives the number of carbon (conjugated)
atoms bound to the metal
It normally, but not necessarily, gives the number of electrons
contributed by the ligand
We will describe to methods of counting electrons but we will
employ only one for the duration of this module
The two methods compared:
some examples
N.B. like oxidation state
assignments, electron
counting is a formalism and
does not necessarily reflect the
distribution of electrons in the
molecule – useful though
Some ligands donate the same
number of electrons
Number of d-electrons and
donation of the other ligands
can differ
Now we will look at practical
examples on the black board
Does it look reasonable ?
• Remember when counting:
• Odd electron counts are rare
• In reactions you nearly always go from even to
even (or odd to odd), and from n to n-2, n or n+2.
• Electrons don’t just “appear” or “disappear”
• The optimal count is 2/8/18 e. 16-e also occurs
frequently, other counts are much more rare.
Exceptions to the 18 Electron Rule
ZrCl2(C5H5)2 Zr(4) + [2 x Cl(1)] + [2 x C5H5(5)] =16
TaCl2Me3 Ta(5) + [2+ x Cl(1)] + [3 x M(1)] =10
WMe6 W(6) + [6 x Me(1)] =12
Pt(PPh3)3 Pt(10) + [3 x PPh3(2)] =16
IrCl(CO)(PPh3)2 Ir(9) + Cl(1) + CO(2) + [2 x PPh3(2)]
=16
What features do these complexes possess?
• Early transition metals (Zr, Ta, W)
• Several bulky ligands (PPh3)
• Square planar d8 e.g. Pt(II), Ir(I)
• σ-donor ligands (Me)