transition complexes
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d-Metal Complexes
Metal complexes consist of a central metal atom or ion surrounded by several
atoms, ions or molecules, calledligands. Ligands are ions or molecules that
can have an independent existence, and are attached to the central metal atomor ion. Examples of ligands are halide ions, carbon monoxide, ammonia,
cyanide ion, etc. In describing complexes, the ligands directly attached to the
metal (usually as Lewis bases, donating electrons to the metal), are counted to
determine the coordination number of the complex. Ions that are directly
coordinated to the metal are written within the brackets of the formula, and are
referred to asinner sphere. Ions that are serving as counter ions in order to
produce a neutral salt, and are not coordinated to the metal are calledouter
sphere, are are written outside of the brackets in the formula.
example: [Mn(OH2)6]SO4: coordination #=6, and sulfate is outer sphere.
[Mn(OH2)5SO4] : coordination # = 6, and sulfate is inner sphere.
Factors Effecting Coordination Number
Coordination numbers can range from 2 up to 12, with 4 and 6 quite common
for the upper transition metals. The following factors influence the coordination
number of the complex.
1.The size of the central atom or ion: Larger atoms (periods 5 & 6) on the
left side of the periodic table are larger, and can accommodate more
ligands.
2. Steric interactions between the ligands: Bulky ligands (such as PPh3) will
result in lower coordination numbers.
3.The electronic structure of the metal: If the oxidation number is high, the
metal can accept more electrons from the Lewis bases. Metals with many
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d electrons, especially those on the right side of the periodic table, will
have lower coordination numbers.
Typical Coordination Numbers and Structures
Coordination
NumberTypical Examples Structure or Shape
2 [CuCl2]-, Ag(NH3)
+ linear
3 (rare) [Cu(CN)2]- trigonal planar
4
CrO42-, NiBr4
2-
favored when metal is small and
ligands are large
tetrahedral
4
[PtCl2(NH3)2], [Ni(CN)4]2-
typical for d8metals and complexes
withbonding ligands
square planar
5
iron porphyrin complexes
(square pyramid due to planar
porphyrin rings)
distorted square
pyramids or trigonal
bipyramids
6
common for d0-d9, most 3d M3+
complexes are octahedral.
distortions include trigonal prism,
especially for chelating ligands with a
small bite angle
7
more common for f-block elements.
Structures include pentagonal
bipyramid, capped octahedron and
capped trigonal prism
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Certain coordination numbers, specifically 5 and 7, have several shapes which
are similar in energy, with low energy barriers for inter conversion. An example
of the fluctuating nature of these shapes can be seen inBerry pseudoratationof
trigonal bipyramidal structures. The structure distorts in such a way that the
axial groups become equatorial, with a square pyramidal structure as an
intermediate.
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Nomenclature of Organometallic Complexes
The system of naming inorganic complexes incorporates prefixes (mono, di, tri,
etc.), to indicate the number of ligands of each type coordinated to the metal.
The ligands have special names (and abbreviations). Typically, ligands which
are negatively charged end ino. A list of common ligands is provided in the text
on page 220. A common ligand which was omitted are phosphines, and
specifically triphenylphosphine, P(C6H5)3. This ligand is often abbreviated as
PPh3or P3. It is important to note that ligands such as this which have
prefixes such as di, tri, etc. in their names are enumerated using the prefixes
bis, tris and tetrakiswhen more than one of them is coordinated to a metal.
The format of naming the complexes is as follows. Ligands are listed first inalphabetical order (not including prefixes). This is followed by the name of the
metal, and the oxidation state of the metal in roman numerals in parenthesis.
prefix for ligands-ligand name-metal name-(oxidation state of metal)
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An alternative system includes the charge on the complex (in arabic numerals)
in parenthesis instead of the oxidation state of the metal. In addition, metal
complexes which are negative in charge often use the latin root for the metal,
and end in the suffixate. The formula of the complex is always written in
square brackets, with the metal appearing first, then negatively charged ions,
and then neutral ligands. Ions which are outer sphere are written outside of
the square brackets.
Name the following compounds:trans-[PtCl2(NH3)4]2+, [Ni(CO)3(py)],
[Cr(edta)]1-
Write formulas for the following compounds:cis-
diaquadichloroplatinum(II), diamminetetra(isothiocyanato)chromium(III),
and tris(ethylenediamine)rhodium(III)
If metals are bridged together by a ligand, the bridging ligand(s) are given the
prefix.
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Multidentate Ligands (Chelating Agents)
Many ligands have more than one site which can coordinate with a metal atom
or ion. These are referred to asambidentateligands. Small ligands with twosites, such as NO2-(N or O) or SCN-(S or N) can undergolinkage isomerism.
The name of the ligand indicates which site is bound to the metal. For NO2-, the
nitrogen linkage has the namenitro, and the oxygen linkage has the name
nitrito. For thiocyanate ion, the sulfur linkage is namedthiocyanato, and the
nitrogen linkage is namedisothiocyanato.
Larger ligands with multiple bonding sites can bond to a single metal atom or
ion at several sites, forming rings. These ligands are referred to aspolydentate,
and formchelate complexes. In the most stable situations, the chelating agent
will form 5 or 6 membered rings with the metal. This produces the properbite
angleto produce octahedral symmetry around the metal. In ligands with smallbite angles, octahedrons will distort to trigonal prismatic geometry.
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One of the highly useful synthetic uses of chelating agents and metal ions
involves thetemplate effect, in which a group of ligands can be assembled by
coordination to a central metal ion. A large variety of macrocyclic compounds
can be synthesized in this way.
Isomerism and Chirality of Metal Complexes
There are several types of isomerism exhibited in transition metal complexes. In
considering only complexes with octahedral geometry around the metal, thefollowing types of isomerism are seen.
1. Geometric Isomerism - The existence ofcisandtransisomers in
compounds with the general formula [MX2Y4]. In thetransisomer, the
two X groups are on opposite sides of the compound (on the same axis).
In thecisisomer, the two X groups occupy adjacent sites. The symmetry
of the molecules vary, and the isomers can be easily detected using
infrared and Raman spectroscopy. Complexes with the general formula
[MX3Y3] can exist as two different isomers. The three like ligands can
either occupy a triangular face of the octahedral structure or three sites
in one plane while the other ligands occupy three sites in a perpendicular
plane. The isomers are designated asfac(for facial) ormer(for
meridional).
2. Chirality and Optical Isomerism - Octahedral complexes can contain
chiral centers, and thus exhibit optical isomerism. The optical isomers,
which are non superimposable mirror images of each other, will bend the
plane of polarized light in different directions. The pair of isomers are
known asenantiomersor an enantiomeric pair. Molecules which arechiral have the absence of an improper axis of rotation. Non-chiral
molecules have a mirror plane through the central atom or a center of
inversion. An example of the formation of many isomers can be seen in
the reaction of cobalt (III) chloride with ethylenediamine, a bidentate
ligand. The products of the reaction include a violet colored product, a
green product, and a yellow product.
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The violet product(s) are the isomers (a) and (b) above. They are mirror images
of each other, and constitute an enantiomeric pair,cis- [CoCl2(en)2]+. The green
product is isomer (c),trans-[CoCl2(en)2]+which lacks optical activity. The yellow
product, [Co(en)3]+
, also exists as an enantiomeric pair. The three bidentateligands can connect to the octahedral sites of the metal in a right handed or
left handed fashion, similar to the blades on a propeller or the threads on a
screw.
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The isomers are designated with the greek symbols(delta, for dextro) for the
right handed complex, and(lambda, for levo) for the left handed complex.
Separation of enantiomers can be accomplished by reaction with a reagent with
a chiral center. This produces molecules which have different solubilities,
melting points, etc.
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Bonding and Electronic Structure of d-Metal Complexes
Crystal Field TheoryThere are two widely used approaches to explaining the bonding and stability of
transition metal complexes. Crystal Field Theory views the electronic field
created by the ligand electron pairs surrounding the central metal as point
negative charges which repel and interact with thedorbitals on the metal ion.
This theory explains the splitting of thedorbitals to remove their degeneracy,
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the number of unpaired electrons in transition metal complexes, their color,
spectra and magnetic properties.
Octahedral Complexes
Examination of the spatial orientation ofdorbitals shows that the dz2and dx2-y2
orbitals point directly at the corners of on octahedron.
The remaining orbitals, the dxy, dxzand dyzare directed between the ligands. As
a result of the octahedral field, the ligands which point between the ligands are
lower in energy, and the orbitals that point directly toward the ligands are
higher in energy. The five degeneratedorbitals are split into two groups.
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The triply degenerate set has the symmetry designation t2g, and the doubly
degenerate set has the designation eg. These designations are from the
character table for the octahedral symmetry group, Oh. The magnitude of the
splitting will depend on the metal and ligands involved. The size of the splitting
is given the symbolo, the ligand field splitting parameter, where the subscript
ostands for octahedral. Since three orbitals are lower in energy and two
orbitals are higher in energy than the original degenerate set, the energy is
lowered by 2/5ofor the t2gset, and raised in energy by 3/5ofor the egset.
The magnitude of the splitting is determined experimentally from spectra of
transition metal complexes. As the size of the gap changes, so does the color ofthe complex, as most of the t2gto egtransitions occur in the visible range.
Analysis of the absorption spectra of a variety of transition metal complexes
has resulted in thespectrochemical series, a list which orders the ligands from
the weakest ligand fields to the strongest.
I-
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1. oincreases with increasing oxidation number of the metal. This is due
to the smaller size of the ion, resulting in smaller metal to ligand
distances, and hence, a greater ligand field.
2. oincreases as you go down a group. This is due to the better bonding
ability of expanded shells using the 4d or 5d orbitals.
The results of these trends is summarized in the list below. The smallest values
ofooccur with the +2 ions, with increasing values observed for higher charged
ions which are lower down in the table.
Mn2+
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The LFSE for the weak field case is equal to [ (3)( 0.40o-(1)(0.60o)] = 0.60o.
The LFSE for the strong field case is equal to (4) (0.40o) = 1.6o. The
different electron configurations are referred to ashigh spin(for the weak field
case) andlow spin(for the strong field case). The possibility of high and low
spin complexes exists for configurations d
5
-d
7
as well. The following generaltrends can be used to predict whether a complex will be high or low spin.
For 3d metals (d4-d7): In general, low spin complexes occur with very
strong ligands, such as cyanide. High spin complexes are common with
ligands which are low in the spectrochemical series, such as the halogen
ions.
For 4d and 5d metals (d4-d7): In general, the size ofois greater than for
3d metals. As a result, complexes are typically low spin. Even a ligand
such as chloride (quite weak) produces a large enough value ofoin thecomplex RuCl6
2-to produce a low spin, t2g4configuration.
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Evidence for LFSE can be seen in the enthalpies of hydration of the 3rdperiod
M2+ions. Without considering LFSE. the enthalpies should increase linearly as
the size of the ion decreases and bond strength increases (the green line).
Instead, the data show two "humps", with the enthalpies for the configurations
d0, d5, and d10falling in the straight line. The other metal ions show greater
enthalpies of hydration (the orange line) in keeping with the calculated ligand-field stabilization energy for each configuration. When the calculated LFSE is
subtracted from the observed values, they fall on the straight line which
ignores LFSE, and only considers ionic size.
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The Electronic Structure of Four-Coordinate Complexes
Tetrahedral and Square Planar Shapes
The d orbitals of tetrahedral complexes also split into two groups. Examination
of the symmetry tables shows that the dxy, dyz, and dxzhave the same symmetry
properties as the px, py, and pzorbitals. If the tetrahedron is viewed as ligandsoccupying the alternating corners of a cube, with the metal atom in the center,
these d orbitals point in the direction of the ligands. As a result, they will be
higher in energy than the degenerate orbitals of the free metal atom or ion. The
dz2and dx2-y2orbitals point between the ligands (towards the center of each
face of the cube), and are lower in energy.
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The splitting pattern has the lower two orbitals (the egset) stabilized by 3/5T,
and the upper three orbitals (the t2set) 2/5Tgreater in energy, whereTis thesize of the splitting in a tetrahedral field. The size ofTis approximately half
the size of the octahedral splittingofor the same metal and ligands, so
virtually all tetrahedral complexes are weak field/high spin.
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Distortions of Octahedral Complexes
Octahedral complexes can undergotetragonal distortions(the elongation of the
z-axis, and shortening of the x and y axes) to become elongated into molecules
with square planar (D4h) symmetry. This distortion is often observed in Cu2+(d9)
octahedral complexes. As thez-axis is elongated, the degeneracy between the
dz2and dx2- y2orbitals is broken, with the dz2orbital lower in energy since the
ligands are further away.
The extreme case of a tetragonal distortion is to form a true square planar
complex in which the ligands along the z-axis are completely removed. In this
case, the dz2orbital drops even lower in energy, and the molecule has the
following orbital splitting diagram.
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As a result of these distortions, there is a net lowering of energy (an increase in
the ligand field stabilization energy) for complexes in which the metal has a d7,
d8
, or d9
configurations, and thus electrons would occupy the upper egset if anoctahedral complex. In general, the size of the splitting in a square planar
complex,SPis 1.3 times greater thanofor complexes with the same metal and
ligands. As a result, the distortion results in square planar complexes with
lower energies than the comparable octahedral complex. This distortion to
square planar complexes is especially prevalent for d8configurations and
elements in the 4thand 5thperiods such as: Rh (I), Ir (I), Pt(II), Pd(III), and Au
(III). Nickel (II) four-coordinate complexes are usually tetrahedral unless there is
a very strong ligand fields such as in [Ni(CN)4]2-, which is square planar.
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The Jahn-Teller Effect
The tetragonal distortions described above are illustrations of theJahn-Teller
Effect. The effect can be summarized by a statement which predicts which
complexes will undergo distortion. It does not predict the extent of the
distortion.
The Jahn-Teller Effect
If the ground electronic configuration of a non-linear complex is orbitally
degenerate, the complex will distort so as to remove the degeneracy and
achieve lower energy.
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This effect is particularly evident in d9configurations. The configuration in a
octahedral complex would be t2g6eg
3, where the configuration has degeneracy
because the ninth electron can occupy either orbital in the egset. A tetragonal
distortion removes the degeneracy, with the electron of highest energy
occupying the non degenerate dx2- y2orbital. Low spin octahedral complexes
with d8configurations are also degenerate, with a square planar distortion
removing any degeneracy.
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Ligand Field Theory
An altnerative approach to understanding the bonding of transition metal
complexes is Ligand Field Theory. Crystal Field Theory is a simple model which
explains the spectra, thermochemical and magnetic data of many complexes.It's main flaw is that it treats the ligands as point charges or dipoles, and fails
to consider the orbitals of the ligands. Ligand Field Theory applies molecular
orbital theory and symmetry concerns to transition metal complexes. In
octahedral symmetry, group theory can be used to determine the shapes and
orientation of the orbitals on the metal and the ligands.
Bonding
Examination of the symmetry table for Oh
shows that the orbitals on the metal
have the following attributes.
metal orbital symmetry label degeneracy
s a1g non-degenerate
px, py, pz t1u triply degenerate
dxy, dyz, dxz t2g triply degenerate
dx2-y2, dz2 eg doubly degenerate
Group theory can be used to determine the combination of ligand orbitals
which have the same symmetry properties as the metal orbitals. The results ofthese Symmetry Adapted Linear Combinations (SALC) are provided below.
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The t2gset (dxy, dyz, dxz) does not have any electron density along the bond axes,
so these orbitals do not participate in sigma bonding, but will be involved with
pi bonding. A molecular orbital diagram which estimates the energies of the
bonding (show above) antibonding and non-bonding orbitals is shown below.
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Since there is a large disparity in energy between the ligand orbitals and the
metal orbitals, the lower lying molecular orbitals in the diagram are essentially
ligand orbitals. That is, the electrons of the ligand lone pairs fill the lower levels
(eg, t1u, and a1g). Thedelectrons on the metal will fill the t2g(non-bonding) and
eg(antibonding) molecular orbitals. The split between the HOMO (highest
occupied molecular orbital) and LUMO (lowest unoccupied molecular orbital)
corresponds to theosplitting in crystal field theory. In crystal field theory, the
electrons in these orbitals are viewed as entirely on the metal atom or ion,
whereas in ligand field theory the electrons are, to some extent, on the ligands,
too.
Bonding
Pi bonding is not considered by crystal field theory, but is addressed in ligand
field theory. The orbitals on the metal which were not used for sigma bonding
(the t2gset: dxy, dyz, dxz) have the same symmetry properties as combinations of
theporbitals on the ligands. If the energy of the metal and ligand orbitals are
comparable, the pi bonding orbitals formed will be significantly lower in energy
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thatn the atomic orbitals on either the metal or ligand. Likewise, the
antibonding pi orbitals will be much higher in energy. If the orbitals are very
different in energy, only slight mixing will occur. An example of pi overlap is
shown below.
The effect on the molecular orbital diagram is as follows. The gap between the
t2gand egset will change, because the t2gset is involved in bonding, so there is
not a bonding t2gset, and an antibonding t2gset of orbitals. The gap,
represented asobecomes the gap between the t2gset of antibonding orbitalsand the egset of orbitals. As a result, the size ofodimishes.
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The above molecular orbital diagram is for ligands which have pi antibonding
orbitals too high in energy to interact with the metal orbitals. The net effect for
these pi donor ligands is to decrease the size ofocompared to ligands which
only act as sigma donors.
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Ligands may have empty pi antibonding orbitals higher in energy and with the
same symmetry as the t2gorbitals of the metal. These ligands orbitals interact
with the t2gorbitals of the metal creating a bonding orbital which is slightly
lower in energy than the t2gset of the metal, and an antibonding set of orbitals
which are much greater in energy than the egset of the complex. The net result
is that the size of the splitting,o, increases, since the energy of the t2gbonding
orbitals drops a bit.
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The net result is that pi acceptor ligands (such as CO and N2), with empty
antibonding orbitals available to accept electrons from the metal, increase the
size ofo. The spectrochemical series can be reconsidered with the possiblity of
pi bonding in mind. It shows that the order (with some notable exceptions) goes
as follows:
strongdonor (smallo) < weakdonor < noeffects (intermediateo)