transition complexes

8/9/2019 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.


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.


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.


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.


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.


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.


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.


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)

transition complexes

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