Download Crystal-Field Theory

Chemistry of Coordination
Compounds
David P. White
University of North Carolina, Wilmington
Chapter 24
Copyright 1999, PRENTICE HALL
Chapter 24
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The Structure of Complexes
• We know Lewis acids are electron pair acceptors.
• Coordination compounds: 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.
• Coordination Sphere: The area of space
encompassing the ligands and metal ion.
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The Structure of Complexes
• Ligands can alter 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
• 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.
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The Structure of Complexes
Charges, Coordination Numbers and Geometries
– Most common coordination numbers are 4 and 6.
– 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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Chelates
• 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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Chelates
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Chelates
[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.
• One very important chelating agent is
ethylenediaminetetraacetate (EDTA4-).
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Chelates
Copyright 1999, PRENTICE HALL
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Chelates
• 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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Chelates
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
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Chelates
Metals and Chelates in Living Systems
Copyright 1999, PRENTICE HALL
Chapter 24
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Chelates
Metals and Chelates in Living Systems
• A five membered nitrogen containing ring binds the
heme unit to the protein.
• When oxygen is attached to the iron(II) in heme,
oxymyoglobin is formed.
• The protein has a molecular weight of about 18,000
amu.
• The Fe2+ ion in oxyhemoglobin or oxymyoglobin is
octahedral.
• Four N atoms from the porphyrin ring (red disk) are
attached to the Fe2+ center.
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Chelates
Metals and Chelates in Living Systems
Copyright 1999, PRENTICE HALL
Chapter 24
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Chelates
Metals and Chelates in Living Systems
• The fifth coordination site is occupied by O2 (or H2O
in deoxyhemoglobin or CO in carboxyhemoglobin).
• The sixth coordination site is occupied by a base,
which attaches the structure to the protein.
• Photosynthesis is the conversion of CO2 and water to
glucose and oxygen in plants in the presence of light.
• One mole of sugar requires 48 moles of photons.
• Chlorophyll absorbs red and blue light and is green in
color.
• The pigments that absorb this light are chlorophylls.
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Chelates
Metals and Chelates in Living Systems
Copyright 1999, PRENTICE HALL
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Chelates
Metals and Chelates in Living Systems
• Chlorophyll a is the most abundant chlorophyll.
• The other chlorophylls differ in the structure of the
side chains.
• Mg2+ is in the center of the porphyrin-like ring.
• The alternating double bonds give chlorophyll its
green color (it absorbs red light).
• Chlorophyll absorbs red light (655 nm) and blue light
(430 nm).
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Chelates
Metals and Chelates in Living Systems
Copyright 1999, PRENTICE HALL
Chapter 24
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Chelates
Metals and Chelates in Living Systems
• The reaction
6CO2 + 6H2O  C6H12O6 + 6O2
is highly endothermic.
• Plant photosynthesis sustains life on Earth.
Copyright 1999, PRENTICE HALL
Chapter 24
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Nomenclature
• 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 tetraamminechlorocobalt(II). Note the
tetra 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).
Copyright 1999, PRENTICE HALL
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Nomenclature
Copyright 1999, PRENTICE HALL
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Nomenclature
• 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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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.
• Stereoisomers have the same connectivity but
different spatial arrangements of atoms.
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Isomerism
Copyright 1999, PRENTICE HALL
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Isomerism
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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Isomerism
Structural Isomerism
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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).
Stereoisomerism
• Consider square planar [Pt(NH3)2Cl2].
• The two NH3 ligands can either be 90 apart or 180
apart.
• Therefore, the spatial arrangement of the atoms is
different.
Copyright 1999, PRENTICE HALL
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Isomerism
Stereoisomerism
Copyright 1999, PRENTICE HALL
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Isomerism
Stereoisomerism
• 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.
• For example, cis-[Co(NH3)4Cl2]+ is violet and trans[Co(NH3)4Cl2]+ is green.
• The two isomers have different solubilities.
Copyright 1999, PRENTICE HALL
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Isomerism
Stereoisomerism
Copyright 1999, PRENTICE HALL
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Isomerism
Stereoisomerism
• 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.)
• 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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Isomerism
Stereoisomerism
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Isomerism
Stereoisomerism
• 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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Isomerism
Stereoisomerism
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Isomerism
Stereoisomerism
• 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 by the same amount in the
opposite direction.
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Isomerism
Stereoisomerism
• Dextrorotatory solutions rotate the plane of polarized
light to the right. This isomer is called the d-isomer.
• Levorotatory solutions rotate the plane of polarized
light to the left. This isomer is called the l-isomer.
• Chiral molecules are optically active because of their
effect on light.
• Racemic mixtures contain equal amounts of l- and disomers. They have no overall effect on the plane of
polarized light.
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Isomerism
Stereoisomerism
• Pasteur was the first to separate racemic ammonium
tartarate (NaNH4C4H9O6) by crystallizing the solution
and physically picking out the “right-handed”
crystals from the mixture using a microscope.
• Optically pure tartarate can be used to separate a
racemic mixture of [Co(en)3]Cl3: if d-tartarate is used,
d-[Co(en)3]Cl3 precipitates leaving l-[Co(en)3]Cl3 in
solution.
Copyright 1999, PRENTICE HALL
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Color and Magnetism
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.
• The color perceived is the sum of the light not
absorbed by the complex.
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Color and Magnetism
Color
Copyright 1999, PRENTICE HALL
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Color and Magnetism
Color
• 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,
– the monochromatic light (i.e. a single wavelength) is passed
through the sample,
– the unabsorbed light is detected.
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Color and Magnetism
Color
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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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Color and Magnetism
Color
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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.
Copyright 1999, PRENTICE HALL
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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.
• The more directly the ligand attacks the metal orbital,
the higher the energy of the d orbital.
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Crystal-Field Theory
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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 dorbitals.
• 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.
• The energy gap between them is called , the crystal
field splitting energy.
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Crystal-Field Theory
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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.
– 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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Crystal-Field Theory
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Crystal-Field Theory
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.
• 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.
•
•
•
•
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Crystal-Field Theory
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Crystal-Field Theory
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Crystal-Field Theory
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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 and the
color of the complex changes from green to yellow.
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Crystal-Field Theory
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Crystal-Field Theory
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 ();
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Crystal-Field Theory
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Crystal-Field Theory
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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.
Copyright 1999, PRENTICE HALL
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Crystal-Field Theory
Tetrahedral and Square-Planar Complexes
• By using the same arguments as for the octahedral
case, we can derive the relative orbital energies for d
orbitals in a tetrahedral field.
• The exact opposite of an octahedral field orbital
arrangement is found (i.e. the dxy, dyz, and dxz orbitals
are of lower energy than the and orbitals).
• Because there are only 4 ligands,  for a tetrahedral
field is smaller than  for an octahedral field (four
ninths).
• This causes all tetrahedral complexes to be high spin.
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Crystal-Field Theory
Copyright 1999, PRENTICE HALL
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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.
• 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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Crystal-Field Theory
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Chemistry of Coordination
Compounds
End of Chapter 24
Copyright 1999, PRENTICE HALL
Chapter 24
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