Metal Complexes
• We know Lewis acids are electron pair acceptors.
• Coordination complexes: metal compounds formed by
Lewis acid-base interactions.
• Complexes: Have a metal ion (can be zero oxidation
state) bonded to a number of ligands. Complex ions are
charged. Example, [Ag(NH3)2]+.
• Ligands are Lewis bases.
• Square brackets enclose the metal ion and ligands.
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Chapter 24
Metal Complexes
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The Development of Coordination
Chemistry: Werner’s Theory
Werner discovered that CoCl3·nNH3 (n = 1 – 4) can exist
as four different compounds with different numbers of
“free” Cl- ions per formula unit.
He deduced that the NH3 ligands were covalently bonded
to the central Co3+ ion.
Werner found that a total of six ligands were attached to
the central Co.
In the case of CoCl3·4NH3, there are two isomers for the
Cl ligands attached to Co.
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Chapter 24
Metal Complexes
The Development of Coordination
Chemistry: Werner’s Theory
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Chapter 24
Metal Complexes
The Metal-Ligand Bond
• All ligands have lone pairs that are donated to the metal
ion.
• The bond between metal and ligand is a 2-electron bond,
but both electrons come from the ligand and none from
the metal.
• The metal-ligand bond alters the physical properties of
the metal:
Ag+(aq) + e-  Ag(s), E = +0.799 V
[Ag(CN)2]-(aq) + e-  Ag(s) + 2CN-(aq), E = -0.031 V
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Chapter 24
Metal Complexes
The Metal-Ligand Bond
• Most metal ions in water exist as [M(H2O)6]n+.
Charges, Coordination Numbers, and Geometries
• Charge on complex ion = charge on metal + charges on
ligands.
• Donor atom: the atom bonded directly to the metal.
• Coordination number: the number of ligands attached to
the metal.
– Most common coordination numbers are 4 and 6.
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Chapter 24
Metal Complexes
Charges, Coordination Numbers, and
Geometries
– Some metal ions have constant coordination number (e.g. Cr3+
and Co3+ have coordination numbers of 6).
– The size of the ligand affects the coordination number (e.g.
[FeF6]3- forms but only [FeCl4]- is stable).
– The amount of charge transferred from ligand to metal affects
coordination number (e.g. [Ni(NH3)6]2+ is stable but only
[Ni(CN)4]2- is stable).
• Four coordinate complexes are either tetrahedral or
square planar (commonly seen for d8 metal ions).
• Six coordinate complexes are octahedral.
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Chapter 24
Ligands with More than
One Donor Atom
• Monodentate ligands bind through one donor atom only.
– Therefore they occupy only one coordination site.
• Polydentate ligands (or chelating agents) bind through
more than one donor atom per ligand.
– Example, ethylenediamine (en), H2NCH2CH2NH2.
• The octahedral [Co(en)3]3+ is a typical en complex.
• Chelate effect: More stable complexes are formed with
chelating agents than the equivalent number of
monodentate ligands.
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Chapter 24
Prentice Hall © 2003
Chapter 24
Ligands with More than
One Donor Atom
[Ni(H2O)6]2+(aq) + 6NH3
[Ni(NH3)6]2+(aq) + 6H2O(l)
Kf = 4  108
[Ni(H2O)6]2+(aq) + 3en
[Ni(en)3]2+(aq) + 6H2O(l)
Kf = 2  1018
• Sequestering agents are chelating agents that are used to
remove unwanted metal ions.
• In medicine sequestering agents are used to selectively
remove toxic metal ions (e.g. Hg2+ and Pb2+) while
leaving biologically important metals.
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Chapter 24
Ligands with More than
One Donor Atom
• One very important chelating agent is
ethylenediaminetetraacetate (EDTA4-).
• EDTA occupies 6 coordination sites, for example
[CoEDTA]- is an octahedral Co3+ complex.
• Both N atoms (blue) and O atoms (red) coordinate to the
metal.
• EDTA is used in consumer products to complex the metal
ions which catalyze decomposition reactions.
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Chapter 24
Prentice Hall © 2003
Chapter 24
Ligands with More than
One Donor Atom
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Metals and Chelates in Living Systems
Many natural chelates are designed around the porphyrin
molecule.
After the two H atoms bound to N are lost, porphyrin is a
tetradentate ligand.
Porphyrins: Metal complexes derived from porphyrin.
Two important porphyrins are heme (Fe2+) and
chlorophyll (Mg2+).
Myoglobin is protein containing a heme unit, which
stores oxygen in cells.
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Chapter 24
Nomenclature of
Coordination Compounds
• Rules:
– For salts, name the cation before the anion. Example in
[Co(NH3)5Cl]Cl2 we name [Co(NH3)5Cl]2+ before Cl-.
– Within a complex ion, the ligands are named (in alphabetical
order) before the metal.
Example [Co(NH3)5Cl]2+ is
pentaamminechlorocobalt(II). Note the penta portion is an
indication of the number of NH3 groups and is therefore not
considered in the alphabetizing of the ligands.
– Anionic ligands end in o and neutral ligands are simply the
name of the molecule. Exceptions: H2O (aqua) and NH3
(ammine).
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Chapter 24
Nomenclature of
Coordination Compounds
• Rules:
– Greek prefixes are used to indicate number of ligands (di-, tri-,
tetra-, penta-, and hexa-). Exception: if the ligand name has a
Greek prefix already. Then enclose the ligand name in
parentheses and use bis-, tris-, tetrakis-, pentakis-, and hexakis.
• Example [Co(en)3]Cl3 is tris(ethylenediamine)cobalt(III) chloride.
– If the complex is an anion, the name ends in -ate.
– Oxidation state of the metal is given in Roman numerals in
parenthesis at the end of the complex name.
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Chapter 24
Isomerism
• Isomers: two compounds with the same formulas but
different arrangements of atoms.
• Coordination-sphere isomers and linkage isomers: have
different structures (i.e. different bonds).
• Geometrical isomers and optical isomers are
stereoisomers (i.e. have the same bonds, but different
spatial arrangements of atoms).
• Structural isomers have different connectivity of atoms.
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Chapter 24
Isomerism
• Stereoisomers have the same connectivity but different
spatial arrangements of atoms.
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Chapter 24
Isomerism
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Chapter 24
Isomerism
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Structural Isomerism
Some ligands can coordinate in different ways.
That is, the ligand can link to the metal in different ways.
These ligands give rise to linkage isomerism.
Example: NO2- can coordinate through N or O (e.g. in
two possible [Co(NH3)5(NO2)]2+ complexes).
When nitrate coordinates through N it is called nitro.
– Pentaamminenitrocobalt(III) is yellow.
• When ONO- coordinates through O it is called nitrito.
– Pentaamminenitritocobalt(III) is red.
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Chapter 24
Isomerism
Structural Isomerism
• Similarly, SCN- can coordinate through S or N.
– Coordination sphere isomerism occurs when ligands from
outside the coordination sphere move inside.
– Example: CrCl3(H2O)6 has three coordination sphere isomers:
[Cr(H2O)6]Cl3 (violet), [Cr(H2O)5Cl]Cl2.H2O (green), and
[Cr(H2O)4Cl2]Cl.2H2O (green).
Structural Isomerism
• Consider square planar [Pt(NH3)2Cl2].
• The two NH3 ligands can either be 90 apart or 180
apart.
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Chapter 24
Isomerism
Stereoisomers
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Chapter 24
Isomerism
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Structural Isomerism
The spatial arrangement of the atoms in the cis and trans
isomers is different.
This is an example of geometrical isomerism.
In the cis isomer, the two NH3 groups are adjacent. The
cis isomer (cisplatin) is used in chemotherapy.
The trans isomer has the two NH3 groups across from
each other.
It is possible to find cis and trans isomers in octahedral
complexes.
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Chapter 24
Isomerism
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Structural Isomerism
For example, cis-[Co(NH3)4Cl2]+ is violet and trans[Co(NH3)4Cl2]+ is green.
The two isomers have different solubilities.
In general, geometrical isomers have different physical
and chemical properties.
It is not possible to form geometrical isomers with
tetrahedra. (All corners of a tetrahedron are identical.)
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Chapter 24
Isomerism
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Structural Isomerism
Optical isomers are mirror images which cannot be
superimposed on each other.
Optical isomers are called enantiomers.
Complexes which can form enantiomers are chiral.
Most of the human body is chiral (the hands, for
example).
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Chapter 24
Isomerism
Structural Isomerism
Prentice Hall © 2003
Chapter 24
Isomerism
Structural Isomerism
Prentice Hall © 2003
Chapter 24
Isomerism
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Structural Isomerism
Enzymes are the most highly chiral substances known.
Most physical and chemical properties of enantiomers are
identical.
Therefore, enantiomers are very difficult to separate.
Enzymes do a very good job of catalyzing the reaction of
only one enantiomer.
Therefore, one enantiomer can produce a specific
physiological effect whereas its mirror image produces a
different effect.
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Chapter 24
Isomerism
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Structural Isomerism
Enantiomers are capable of rotating the plane of
polarized light.
Hence, they are called optical isomers.
When horizontally polarized light enters an optically
active solution.
As the light emerges from the solution, the plane of
polarity has changed.
The mirror image of an enantiomer will rotate the plane
of polarized light in the opposite direction.
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Chapter 24
Color and Magnetism
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Color
Color of a complex depends on: (i) the metal and (ii) its
oxidation state.
Pale blue [Cu(H2O)6]2+ can be converted into dark blue
[Cu(NH3)6]2+ by adding NH3(aq).
A partially filled d orbital is usually required for a
complex to be colored.
So, d0 metal ions are usually colorless. Exceptions:
MnO4- and CrO42-.
Colored compounds absorb visible light.
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Chapter 24
Color and Magnetism
Color
• The color perceived is the sum of the light not absorbed
by the complex.
• The amount of absorbed light versus wavelength is an
absorption spectrum for a complex.
• To determine the absorption spectrum of a complex:
– a narrow beam of light is passed through a prism (which
separates the light into different wavelengths),
– the prism is rotated so that different wavelengths of light are
produced as a function of time,
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Chapter 24
Color and Magnetism
Color
– the monochromatic light (i.e. a single wavelength) is passed
through the sample,
– the unabsorbed light is detected.
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Chapter 24
Color and Magnetism
Color
• The plot of absorbance versus wavelength is the
absorption spectrum.
• For example, the absorption spectrum for [Ti(H2O)6]3+
has a maximum absorption occurs at 510 nm (green and
yellow).
• So, the complex transmits all light except green and
yellow.
• Therefore, the complex is purple.
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Chapter 24
Prentice Hall © 2003
Chapter 24
Color and Magnetism
Magnetism
• Many transition metal complexes are paramagnetic (i.e.
they have unpaired electrons).
• There are some interesting observations. Consider a d6
metal ion:
– [Co(NH3)6]3+ has no unpaired electrons, but [CoF6]3- has four
unpaired electrons per ion.
• We need to develop a bonding theory to account for both
color and magnetism in transition metal complexes.
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Chapter 24
Crystal-Field Theory
• Crystal field theory describes bonding in transition metal
complexes.
• The formation of a complex is a Lewis acid-base
reaction.
• Both electrons in the bond come from the ligand and are
donated into an empty, hybridized orbital on the metal.
• Charge is donated from the ligand to the metal.
• Assumption in crystal field theory: the interaction
between ligand and metal is electrostatic.
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Chapter 24
Crystal-Field Theory
• The more directly the ligand attacks the metal orbital, the
higher the energy of the d orbital.
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Chapter 24
Crystal-Field Theory
• The complex metal ion has a lower energy than the
separated metal and ligands.
• However, there are some ligand-d-electron repulsions
which occur since the metal has partially filled d-orbitals.
• In an octahedral field, the degeneracy of the five d
orbitals is lifted.
• In an octahedral field, the five d orbitals do not have the
same energy: three degenerate orbitals are higher energy
than two degenerate orbitals.
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Chapter 24
Crystal-Field Theory
• The energy gap between them is called , the crystal field
splitting energy.
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Chapter 24
Crystal-Field Theory
• We assume an octahedral array of negative charges
placed around the metal ion (which is positive).
• The and orbitals lie on the same axes as negative
charges.
– Therefore, there is a large, unfavorable interaction between
ligand (-) and these orbitals.
– These orbitals form the degenerate high energy pair of energy
levels.
• The dxy, dyz, and dxz orbitals bisect the negative charges.
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Chapter 24
Crystal-Field Theory
− Therefore, there is a smaller repulsion between ligand and metal
for these orbitals.
– These orbitals form the degenerate low energy set of energy
levels.
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The energy gap is the crystal field splitting energy .
Ti3+ is a d1 metal ion.
Therefore, the one electron is in a low energy orbital.
For Ti3+, the gap between energy levels,  is of the order
of the wavelength of visible light.
Prentice Hall © 2003
Chapter 24
Crystal-Field Theory
• As the [Ti(H2O)6]3+ complex absorbs visible light, the
electron is promoted to a higher energy level.
• Since there is only one d electron there is only one
possible absorption line for this molecule.
• Color of a complex depends on the magnitude of 
which, in turn, depends on the metal and the types of
ligands.
Prentice Hall © 2003
Chapter 24
Crystal-Field Theory
• Spectrochemical series is a listing of ligands in order of
increasing :
Cl- < F- < H2O < NH3 < en < NO2- (N-bonded) < CN• Weak field ligands lie on the low end of the
spectrochemical series.
• Strong field ligands lie on the high end of the
spectrochemical series.
• As Cr3+ goes from being attached to a weak field ligand
to a strong field ligand,  increases.
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Chapter 24
Prentice Hall © 2003
Chapter 24
Crystal-Field Theory
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Electron Configurations in Octahedral
Complexes
We still apply Hund’s rule to the d-orbitals.
The first three electrons go into different d orbitals with
their spins parallel.
Recall: the s electrons are lost first.
So, Ti3+ is a d1 ion, V3+ is a d2 ion and Cr3+ is a d3 ion.
We have a choice for the placement of the fourth
electron:
– if it goes into a higher energy orbital, then there is an energy
cost ();
Prentice Hall © 2003
Chapter 24
Crystal-Field Theory
Electron Configurations in Octahedral
Complexes
– if it goes into a lower energy orbital, there is a different energy
cost (called the spin-pairing energy due to pairing an electron).
• Weak field ligands tend to favor adding electrons to the
higher energy orbitals (high spin complexes) because  <
pairing energy.
• Strong field ligands tend to favor adding electrons to
lower energy orbitals (low spin complexes) because  >
pairing energy.
Prentice Hall © 2003
Chapter 24
Crystal-Field Theory
Tetrahedral and Square-Planar
Complexes
• Square planar complexes can be thought of as follows:
start with an octahedral complex and remove two ligands
along the z-axis.
• As a consequence the four planar ligands are drawn in
towards the metal.
• Relative to the octahedral field, the orbital is greatly
lowered in energy, the dyz, and dxz orbitals lowered in
energy, the dxy, and orbitals are raised in energy.
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Chapter 24
Crystal-Field Theory
Tetrahedral and Square-Planar
Complexes
• Most d8 metal ions for square planar complexes.
– the majority of complexes are low spin (i.e. diamagnetic).
– Examples: Pd2+, Pt2+, Ir+, and Au3+.
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Chapter 24
Tetrahedral and
Square-Planar
Complexes
• Most d8 metal ions for
square planar
complexes.
– the majority of
complexes are low spin
(i.e. diamagnetic).
– Examples: Pd2+, Pt2+,
Ir+, and Au3+.