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Transition metal complexes
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A transition metal complex is species
consisting of a transition metal coordinated
(bonded to) one or more ligands (neutral or
anionic non-metal species)
Transition metal complexes are important
in catalysis, materials synthesis,
photochemistry, and biological systems
Display diverse chemical, optical and
magnetic properties
Transition metals
Coordination numbers
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Transition metal ions usually form
complexes with a well defined number of
ligands
Complexes with coordination numbers four
and six are the most common, although five
coordination is also well established
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Stereochemistry
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Four coordinate transition metal ions adopt
either tetrahedral or square planar
coordination
– FeCl42- (tetrahedral), AuCl4- (square planar)
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Six coordinate species are nearly always
octahedral
Five coordinate species are either trigonal
bipyramidal or square base pyramidal
Example geometries
Ligands
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Ligands are species (neutral or anionic)
bonded to the metal ion
They may be attached to the metal through
a single atom (monodentate) or bound to
the metal through two or more atoms
(bidentate, tridentate etc.)
Polydentate ligands are called chelating
ligands
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Example chelating ligands
Ligands and oxidation state
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Low oxidation state complexes can be
stabilized by using ligands such as cyanide
and carbon monoxide
Intermediate oxidation state complexes
often have ligands such as chloride,
ammonia or water
High oxidation state complexes usually
have fluoride or oxide as ligands
Isomerism in metal complexes
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Geometrical isomers
Optical isomerism
Linkage isomerism
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Some ligand can link to a metal through
one of two or more different atoms
– NO2- can link through oxygen or nitrogen
» [Co(ONO)(NH 3)5]+ and [Co(NO 2)(NH 3)5]+
– NCS- can link through S or N
– SO32- can link through O or S
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Ionization and hydration isomerism
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Ionization isomerism – [Co(NH3)5(Br)]SO4 versus [Co(NH3)5(SO4)]Br
– or [Pt(NH3)3(Cl)]Br versus [Pt(NH3)3(Br)]Cl
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Hydration isomerism
– [Cr(H2O)6]Cl 3 versus [Cr(H 2O)5Cl]Cl 2.H 2O versus
[Cr(H2O)4Cl 2]Cl.2H 2O
Naming of complexes
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Your own responsibility
Should be able to name a variety of simple
coordination complexes
Bonding in metal complexes
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A number of theories are commonly used
18 electron rule - only applies to low
oxidation state complexes, particularly
organometallics (things with carbon bonded
to the metal center)
Crystal field theory - an electrostatic model
– pretty good for many complexes
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MO theory
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18 electron rule
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Stable low oxidation state complexes are
found to have a total of 18 bonding
electrons
– metal electrons plus lone pairs from ligands
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Ni(CO)4 - 4s23d8 and 4 lone pairs
Fe(CO)5 - 4s23d6 and 5 lone pairs
Cr(CO)6 - 4s23d4 and 6 lone pairs
Can be explained using MO theory
Crystal Field Theory
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Consider the ligands are point negative
charges or as dipoles. How do these charges
interact with the electrons in the d-orbitals?
Octahedral complexes
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Ligand field splitting
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The ligand field splitting depends upon the
oxidation state of the metal and the ligand
type
High oxidation state favors large D
Effect of ligand is given by the
spectrochemical series
– I-< Br-< Cl -< F-< OH-< OH 2< NH3< en< CN-< CO
High spin and low spin complexes
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HS versus LS is determined by the relative
size of the ligand field splitting and the
pairing energy
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If Dis bigger than the pairing energy the
complex will be low spin
Tetrahedral complexes
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Square planar complexes
Magnetic properties
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The loss of degeneracy of the d-orbitals due
to crystal field splitting explains why some
complexes are diamagnetic and others are
paramagnetic
– e.g. Ni(CN)42- (square planar) is diamagnetic
– but NiCl42- (tetrahedral) is paramagnetic
Colors
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The colors of most transition metal
complexes arises as a consequence of the dorbital splitting
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Hydration energies
Electronic spectra
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TM complexes with more than one delectron often show absorption bands at
multiple wavelengths
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The appearance of multiple bands is due to
electron-electron repulsion
– not just one transition corresponding to D
– may have more than one state for a given
electron configuration
d2 complexes
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MO theory
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Crystal field theory explains many
phenomena but not for instance the
spectrochemical series
MO theory offers a better qualitative
picture
Octahedral complex and sigma donor ligands
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Dis the energy gap between nonbonding
orbitals and antibonding orbitals
p -donor ligands
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p -donor ligands(e.g. halide)give smaller
Ds
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p -acceptor ligands
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p -acceptor ligands (e.g. CO) give large Ds
Carbonyl complexes
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Carbonyl ligands are strong p -acceptors
– backbonding weakens the C=O bond as seen by
vibrational spectroscopy
Coordination equilibria
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For many complexes in solution there is an
equilibrium between different species
– Ni(OH2)62+ + 6NH3 = Ni(NH3)62+ + 6H2O
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Equilibrium constant depends upon metal
and ligands
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Chelate effect
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The equilibrium constants for chelate
complexes are usually larger than for
similar non-chelate complexes
– Ni(OH2)62+ + 6NH3 = Ni(NH3)62+ + 6H2O K1
– Ni(OH2)62+ + 3en = Ni(en)32+ + 6H2O K 2
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K2 > K1
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The effect is due to entropic differences
between the two reactions
Kinetics
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For many complexes the exchange of
ligands in the coordination sphere occurs
rapidly
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However, for complexes of Co(III) and
Cr(III) ligand exchange kinetics are slow
– labile
– said to be kinetically inert
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