The document discusses valence bond theory and crystal field theory as they apply to coordination compounds. Valence bond theory describes chemical bonding as occurring through the overlap of atomic orbitals. It can explain octahedral complexes through d2sp3 or sp3d2 hybridization of the metal's orbitals. Crystal field theory postulates that ligands exert an electrostatic field that splits the metal's d-orbitals into two energy levels. In an octahedral field, the eg orbitals have higher energy than the t2g orbitals. Complexes with weak ligands that cause little splitting tend to be high spin, while those with strong ligands that cause large splitting are low spin.

1. COORDINATION COMPUNDS
 VALENCE BOND THEORY
 CRYSTAL FIELD THEORY

2. VALENCE BOND THEORY
Linus Pauling
in
1962

3. G.N. LEWIS
1916
WLTER HEITLE AND
FRITZ LONDON
Heitler- Londin theory
1927
Enables to calculate
nonbonding properties
of covalently bonded
hydroden
Chemical bonding is in the
ability of atoms to share two
bonding electrons
Valence
bond theory
Linus Pauling
RESONANCE
&
ORBITAL HYBRIDISATION

4.  It is a quantum mechanical model to explain the chemical bonding formation.
 This theory primarily focuses on the formation of individual bonds from the atomic
orbitals of the participating atoms during the formation of a molecule.
WHAT IS VB THEORY
According to this theory,
 Electrons in a molecule occupy atomic orbitals rather than molecular orbitals.
 The atomic orbitals overlap on the bond formation.
ORBITAL OVERLAP an orbital on one atom comes to occupy a portion of
the same region of space as an orbital on the other atom.
COVALENT BOND formed between two atoms by the overlap of half filled
valence atomic orbitals of each atom
containing one unpaired electron.
LARGER THE OVERLAP STRONGER THE BOND STRENGTH

5. ASSUMPTION OF VBT:
 The theory assumes that electrons occupy atomic orbitals of individual atoms
within a molecule, and that the electrons of one atom are attracted to the nucleus of
another atom.
 This attraction increases as the atoms approach one another until the atoms
reach a minimum distance where the electron density begins to cause repulsion
between the two atoms.
This electron density at the minimum distance between the two atoms is where
the lowest potential energy is acquired, and it can be considered to be what holds
the two atoms together in a chemical bond.
COVALENT BOND STRUCTURE LEWIS STRUCTURE
TWO TYPES OF BONDING:
Sigma bond formation pi bond formation

6. HYBRIDISATION Atomic orbitals that are similar in energy
but not
equivalent are combined mathematically to
produce equivalent orbitals that are properly
oriented to form bonds.
COMBINATION OF
TWO OR MORE new combinations
ATOMIC ORBITALS
OF HYBRID ATOMIC ORBITALS
SAME ATOM

7. Hybdsation pic

8. Vbt in coordination compounds
VALENCE BOND THEORY
IN
COORDINATION COMPOUNDS

9. COORDINATION COMPOUNDS METAL - LIGAND
COMBINATION
The metal bonding is essentially covalent in origin
 Metallic structure involves resonance of electron-pair bonds between each
atom and its neighbours.
 This VBT deals with the electronic structure of central metal ion in its ground
state, kind of bonding geometry and magnetic properties of the complex.
 It is easy to apply the valence-bond theory to some coordination complexes,
such as the Co(NH3)6
3+ ion.

10. electron configuration of the transition- metal ion.
Co3+: [Ar] 3d6
We then look at the valence-shell orbitals and note that the 4s and 4p orbitals are empty.
Co3+: [Ar] 3d6 4s0 4p0
Concentrating the 3d electrons in the dxy, dxz, and dyz orbitals in this subshell gives the
following electron configuration.
3dx
2
-y
2, 3dz
2, 4s, 4px, 4py and 4pz orbitals are then mixed to form a set of empty d2sp3 orbitals
that point toward the corners of an octahedron
 Each of these orbitals can accept a pair of nonbonding electrons from a neutral
NH3 molecule to form a complex in which the cobalt atom has a filled shell of valence
electrons.

12. Postulates
POSTULATES
OF
VALENCE BOND THEORY

13.  The metal atom or ion in the complex makes available no of empty orbitals,
which can accommodate electrons donated by the ligands to form coordination
compounds.
 No of empty orbitals Coordination of central metal ion
 The appropriate atomic orbitals of metal orbitals hybridize to give an equal no
of new orbitals of equivalent energy called hybrid orbital.
 The hybrid orbitals are directed towards the ligand positions according to
the geometry of the complex.
Atoms retain their individuality in the molecule
 The bond is formed due to the interaction of valence electrons as the atoms
approach each other i.e., inner orbitals from each atoms forming the bond are
undisturbed

14. The d-orbitals involved in hybridisation may be inner viz., (n-1)d orbitals or
outer viz., nd orbitals.
Eg., Octahedral hybridisation may be either
(n-1)d 2sp3 Low spin complex
Or nsp3d2 High spin complex
Each ligand has at least one orbital containing a lone pair of electrons.
 The empty hybrid orbitals of metal ion overlap with fully filled orbitals of the
ligand forming the ligand – metal coordinate bond
 Extent of overlapping Stability of the complex
If complex contains,
Unpaired electrons Paramagnetic complex
Paired electrons Diamagnetic complex
 The non bonding electrons present in the inner orbitals do not take part in
chemical bonding.

15. OCTAHEDRAL COMPLEX
INNER ORBITAL
COMPLEX
OUTER ORBITAL COMPLEX
D2SP3
HYBRIDISATION
Sp3d2
HYBRIDISATION
This hybridisation is depends upon the paired and unpaired electrons
present in the complex
Central
atom
COORDINATION
NUMBER : 6

16. INNER ORBITAL OCTAHEDRAL COMPLEXES
 The octahedral complex compounds resulted from d2sp3 hybridisation are called inner
orbital octahedral complexes.
It takes place in those octahedral complexes which contain strong ligands.
On the basis of orientation of lobes of d orbitals in space, these orbitals have been
classified into two sets viz. t2g and eg
t2g dxy, dyz, dzx orbitals
eg dz2 , dx2-y2 orbitals
 In the formation of six d2sp3 hybrid orbitals,
Two d-orbitals eg set (n-1) dz2 & (n-1)dx2-y2 INNER SHELL
Both are from penultimate (n-1) shell
Low spin complexes having lesser number of
OR unpaired electrons due to paring
Spin paired octahedral complexes
1 2 3 4 5 6
3d 4s 4p
D2sp3 hybridisation

17.  Co3+ does not contain any empty 3d orbital . Now according to VBT, in the
presence of CN- (strong ligand) the electrons in 3d orbitals are force to pair up
against Hund’s rule in order to make room for electrons donated by the ligands.
 Now there are six hybrid d2sp3 orbitals which can accept six electron pair form
ligands.
d2sp3 hybridization gives octahedral geometry
 It has no unpaired electrons, so it is diamagnetic in nature.
 The inner (n-1)d orbitals are taking part in bond formation, hence it is inner
complex ion
Eg, [Co(CN)6]3- ,[ Co(No2)] 3-

18. OUTER ORBITAL OCTAHEDRAL COMPLEX
 The octahedral complexes in which the central atom is sp3d2 hybridized are called outer
orbital complex
This type of hybridisation takes place in those octahedral complex ions which contain weak
ligands
 Weak ligands are those which cannot force the electrons of dz2 and dx2-y2 orbitals (eg set)
of the inner shell.
Thus in this hybridisation (n-1)dz2 and (n-1)dx2-y2 orbitals are not available for hybridisation.
 Orbitals of outer shell ( ns, npx npy, npz, ndxy,ndx2-y2)
are used in this hybridisation.
Since two d-orbitals from outer shell (nth shell), the octahedral complexes resulted form
sp3d2hybridisation are called outer orbital octahedral complexes.
1 2 3 4 5 6
3d 4s 4p 4d

19. High spin Greater number of
OR unpaired electrons
spin free octahedral complex than inner orbital complex
 This complex is found to be paramagnetic since there are four unpaired electrons.
 This is because the weak nature of the ligand
 It cannot force the pairing up of electrons against Hund’s rule.
 The outer nd orbitals taking part in bond formation, hence it is outer complex
ion.

20. LIMITATIONS OF VALENCE BOND THERORY
 The basic defect in the theory is that attention is focused only on the orbitals
of central atom & the importance of the ligands are not stressed.
It does not explain the relative energies of different structures
It does not explain the spectra of complexes
It offers no explanation of the colour observed for complex ion
It does not take account the splitting of d-energy levels
It cannot predict whether a 4-coordination complex will be tetrahedral or square
planar.
The VBT does not provide any satisfactory explanation for the existence of
inner &outer orbital complexes.
It cannot explain the inertness and labiality of the complexes
This theory does not predict or explain the magnetic behavior of complexes
and it predicts only the no of unpaired electrons.

21. INNER ORBITAL COMPLEX OUTER ORBITAL COMPLEX
Central atom is d2sp3 hybridisation Central atom is sp3d2 hybridisation
(n-1)d orbital involved here is belong to
inner
(penultimate )shell while ns & np are belong
to outer (ultimate)shell
Here all the orbitals belong to the ultimate
(outer) shell
These complexes are given by strong ligands These complexes are given by weak ligands
Generally have paired electrons Generally have unpaired electrons
Diamagnetic in nature Paramagnetic in nature
Spin paired or low spin or hyper ligated
complex
Spin free or high spin or hypo ligated
complexes

22. cft
CRYSTAL FIELD THEORY

23. Crystal field theory (CFT) describes the breaking of orbital degeneracy
in transition metal complexes due to the presence of ligands.
 CFT qualitatively describes the strength of the metal-ligand bonds.
POSTULATES OF CFT
 The ligands are point charges which are either ions [ F-, Cl-, CN-, etc] or neutral
molecules [NH3, H2O,etc] with negative poles oriented towards metal cation
The transition metal cation is surrounded by the ligand with lone pair of
electrons
The attraction between metal cation and ligands is purely ionic
The valence electrons of metal will be repelled by negative field of ligands so these
electrons occupy those d-orbitals which have their lobe away from direction of the
ligand
 The colour of transition metal complexes can be explained in terms of electronic
transition between the various d-orbitals of different energy
The magnetic properties can be explained in terms of splitting of d-orbitals in
different crystal field.
Different crystal fields will have different effect of the relative energies of the five
d-orbitals.

24. CRYSTAL FIELD SPLITTING
Splitting of five degenerate d-orbitals of metal ion under the influence of
approaching ligands into two sets of orbital having different energies.
The energy gap between the two sets of orbitals is called
crystal field splitting energy or crystal field stabilization energy.
Fig: splitting of degenerated d-orbitals

25. SPLITTING OF d-ORBITALS IN OCTAHEDRAL FIELD
d-ORBITAL
eg t2g
Lobes are along the axis
 They are called axial orbitals
 dz2 and dx2-y2
“e” doubly degenerate
set
 lobes are between the axis
They are called non-axial orbitals
dxy, dxz, dyz
“t” triply degenerate set

27.  In an octahedral complex, the metal ion is at the centre of the octahedron
and the ligand are at six corners of the octahedron.
 In case of free metal ion all the five d-orbital are degenerate i.e. they all
have same energy.
 On the approach of ligands (along the axis) the electrons in the d-orbitals of
central metal ion are repelled by the lone pair of ligands.
 The repulsion will raise the energy of all the five d-orbitals
 Since, the lobes of the two eg orbitals lie directly in the path of the
approaching ligands, the electrons in these orbitals experience greater force of
repulsion than in t2g orbitals whose lobes are directed in space between the path
of approaching ligands.
 i.e., Energy of eg orbitals increased while that of t2g is decreased.

28. Greater repulsion Greater energy
t2g orbital lower energy
eg orbital Higher energy
 The energy gap between t2g and eg sets is denoted by ∆0 or 10Dq
where, subscript 0 stands for octahedral.
 The magnitude of splitting is arbitrarily fixed as 10Dq so that ∆o = 10Dq
 t2g orbitals are 4Dq less than the degenerate level.
 eg orbitals are 6Dq greater than that of degenerate level.
 Each electron entering the t2g orbitals stabilizes the complex ion by 4Dq
Each electron entering the higher energy, eg orbitals destabilizes by an amount of 6Dq.
STRONG FIELD LIGANDS Ligands produce greater crystal field splitting energy
WEAK FIELD LIGANDS Ligands produce lesser crystal field splitting energy
On the basis of crystal field splitting energy, some common ligands in increasing
order of this field strength are listed below
It is called SPECTRO CHEMICAL SERIES
I– < Br– < Cl– < SCN– < F– < OH– < C2O4
2- < H2O < NCS– < EDTA4- < NH3 < en < CN–<
CO

29. HIGH SPIN AND LOW SPIN COMPLEXES
The complexion with the greater number of unpaired electrons is known as the high spin
complex
The low spin complex contains the lesser number of unpaired electrons.
 High spin complexes are expected with weak field ligands whereas the crystal field
splitting energy is small Δ.
 The opposite applies to the low spin complexes in which strong field ligands cause
maximum pairing of electrons in the set of three t2 atomic orbitals due to large Δo.
High spin Maximum number of unpaired electrons.
Low spin Minimum number of unpaired electrons
[Co(CN)6]3- – Low spin complex
[CoF6]3- – High spin complex

30. CRYSTAL FIELD STABILIZATION ENERGY
The energy of the electron configuration in the ligand field minus the energy of the
electronic configuration in the isotropic field.
CFSE = ∆E = Eligand field - Eisotopic field .
 The energy difference between the eg and t2g levels is given as ∆ or 10Dq
 It states that each electron that goes into the lower t2g level stabilizes the system by an
amount of -4Dq and the electron that goes into eg level destabilizes the system by +6Dq.
 That is the t2g is lowered by 4Dq and the eg level is raised by +6Dq.
Paring in
Degenerate
d - orbital
[ (t2g electrons) × (- 4Dq) ] + [ (eg electrons) × ( 6Dq) + paring energy
For example, the net change in energy for d5 and d10 systems will be zero as
shown below.
d5 :- 3(-4Dq) + 2(+6Dq) = -12Dq + 12Dq = 0
d10 :- 6(-4Dq) + 4(+6Dq) = -24Dq + 24Dq = 0

32. SPLITTING IN TETRAHEDRAL FIELD
 The splitting of fivefold degenerate d orbitals of the metal ion into two levels in a
tetrahedral crystal field is the representation of two sets of orbitals as Td
Two set t2 & e
 In tetrahedral complex, there are four ligands attached to the central metal atom
along the axis.
 The ligands will approach the orbitals along the axis more than that orbitals in
between the axis.

33. The electrons in dx
2
-y
2 and dz
2 orbitals are less repelled by the ligands than
the electrons present in dxy, dyz, and dxz orbitals.
 As a result, the energy of dxy, dyz, and dxz orbital set are raised while that the
energy of dx
2
-y
2 and dz
2 orbitals are lowered.
 Thus, the repulsions in tetrahedral coordination compound yield two energy
levels:
t2– set of three orbitals (dxy, dyz and dxz) with higher energy
e – set of two orbitals (dx
2
-y
2 and dz
2) with lower energy
The crystal field splitting in a tetrahedral complex is intrinsically smaller in
an octahedral filed because there are only two thirds as many ligands and they
have a less direct effect of the d orbitals.
 The relative stabilizing effect of e set will be -6Dq and the destabilizing
effect of t2 set will be +4Dq

34. SPLITTING IN SQUARE PLANAR COMPLEX
 In square planar complex, there are four ligands along the x and y axis
The electrons of the ligands are only attracted to the xy plane.
 Any orbital in the xy plane has a higher energy level
 As ligands approach through the axis they would have greater influence on
dx2-y2 orbital so that energy of this orbital will be most
The dxy orbital lying in the same plane but between the ligands will also
have greater energy though the effect will be less than that on dx2-y2 orbitals.
 On other hand due to absence of ligand in z axis , dz2 orbitals becomes
stable and has lower energy than dxy orbital s
 similarly dyz, dxz become more stable.

37. Factors affecting
FACTORS
AFFECTING
CRYSTAL FIELD THEORY

38. CHARGE ON THE METAL ION:
Increasing the charge on a metal ion has two effects:
the radius of the metal ion decreases,
and negatively charged ligands are more strongly attracted to it.
 Both factors decrease the metal–ligand distance, which in turn causes the
negatively charged ligands to interact more strongly with the d orbitals.
 Consequently, the magnitude of Δo increases as the charge on the metal ion
increases.
 Typically, Δo for a tripositive ion is about 50% greater than for the dipositive ion
of the same metal;
for example,
for [V(H2O)6]2+, Δo = 11,800 cm−1;
for [V(H2O)6]3+, Δo = 17,850 cm−1.
PRINCIPAL QUANTUM NUMBER OF THE METAL:
For a series of complexes of metals from the same group in the periodic table
with the same charge and the same ligands, the magnitude of Δo increases with
increasing principal quantum number: Δo (3d) < Δo (4d) < Δo (5d).

39. The data for hexamine complexes of the trivalent Group 9 metals illustrate this
point:
[Co(NH3)6]3+: Δo = 22,900 cm−1
[Rh(NH3)6]3+: Δo = 34,100 cm−1
[Ir(NH3)6]3+: Δo = 40,000 cm−1
The increase in Δo with increasing principal quantum number is due to the larger
radius of valence orbitals down a column.
 In addition, repulsive ligand–ligand interactions are most important for
smaller metal ions.
This results in shorter M–L distances and stronger d orbital–ligand interactions
NATURE OF THE LIGAND:
 Ligands which cause only a small splitting are called weak ligands and those
which bring a large splitting is called strong ligands.
It is possible to arrange the ligands according to magnitude of ∆ observed
with metal ion.

40.  The sequence for common ligands in order of increasing ∆values is as follow
OXIDATION STATE OF THE METAL ION:
 The magnitude of ∆ increases as the oxidation state of the metal ion increases
The splitting produced by +3 oxidation state is roughly 1.5 times or doubly than
that of +2 oxidation state.
GEOMETRY OF THE COMPLEX:
 The magnitude of ∆ varies with the geometry of complex
Eg., ∆ values for tetrahedral complexes is about 40-5-% less than that of
octahedral complexes for the same ligand
 It has been found that ∆ for square planar complexes is more than that for
octahedral complexes.

41. apps
APPLICATIONS
OF
CRYSTAL FIELD THEORY

42. COLOUR OF THE TRANSITION METAL COMPLEXES
 CFT provides a satisfactory explanation for the observed colour in the
complexes
Absorption of light visible region of electromagnetic spectrum
By coloured complex 4000A to 7000A
 because ∆ 0 value for most of the complexes is of the same order as the energy of
visible radiation.
 The striking colours exhibited by transition-metal complexes are caused by
excitation of an electron from a lower-energy d orbital to a higher-energy d
orbital, which is called a d–d transition
 The absorbed visible or ultra violet absorption spectrum corresponds to
promotion of an electron form t2g orbital of an eg orbital in octahedral complex
The absorbed light is different from that of the transmitted light

43. For example,
The complex [Cr(NH3)6]3+ has strong-field ligands and a relatively large Δo.
Consequently, it absorbs relatively high-energy photons, corresponding to
blue-violet light, which gives it a yellow color.
 A related complex with weak-field ligands, the [Cr(H2O)6]3+ ion, absorbs
lower-energy photons corresponding to the yellow-green portion of the visible
spectrum, giving it a deep violet color

44. It is clear that the environment of the transition-metal ion, which is determined by
the host lattice, dramatically affects the spectroscopic properties of a metal ion.
Fig: Gem-quality crystals of ruby and emerald. The colours of both minerals are due to
the presence of small amounts of Cr3+ impurities in octahedral sites in an otherwise
colourless metal oxide lattice.
RUBY :
 Cr–O distances are relatively short because of the constraints of the host lattice, which
increases the d orbital–ligand interactions and makes Δo relatively large.
Consequently, rubies absorb green light and the transmitted or reflected light is red,
which gives the gem its characteristic color
EMERALD :
Cr–O distances are longer due to relatively large [Si6O18]12− silicate rings; this results in
decreased d orbital–ligand interactions and a smaller Δo.
Consequently, emeralds absorb light of a longer wavelength (red), which gives the gem
its characteristic green color

45. MAGNETIC PROPERTIES OF COORDINATION COMPLEX
An interesting characteristic of transition metals is their ability to form magnets.
Metal complexes that have unpaired electrons are magnetic.
Since the last electrons reside in the d orbitals, this magnetism must be due to having
unpaired d electrons.
 The magnetic properties of a compound can be determined from its electron
configuration and the size of its atoms.
 Because magnetism is generated by electronic spin, the number of unpaired
electrons in a specific compound
PARAMAGNETISM:
Paramagnetism refers to the magnetic state of an atom with
one or more unpaired electrons.
The unpaired electrons are attracted by a magnetic field due to the electrons' magnetic
dipole moments.
 Hund's Rule states that electrons must occupy every orbital singly before any
orbital is doubly occupied.
This may leave the atom with many unpaired electrons. Because unpaired electrons can
spin in either direction, they display magnetic moments in any direction.
 This capability allows paramagnetic atoms to be attracted to magnetic fields.

46. The size of the magnetic moment of a system containing unpaired electrons is related
directly to the number of such electrons
The greater the number of unpaired electrons, the larger the magnetic moment.
 Magnetic susceptibility measures the force experienced by a substance in a magnetic
field.
 When we compare the weight of a sample to the weight measured in a magnetic
field , paramagnetic samples that are attracted to the magnet will appear heavier
because of the force exerted by the magnetic field.
We can calculate the number of unpaired electrons based on the increase in weigh
From this experiment, the measured magnetic moment of low-
spin d6 [Fe(CN)6]4− ion confirms that iron is diamagnetic, whereas the high-
spin d6 [Fe(H2O)6]2+ complex has four unpaired electrons with a magnetic moment
that confirms this arrangement.
Gouy’s Balance
method

47. DIAMAGNETISM:
 Diamagnetic substances are characterized by paired electrons i.e., there are
no unpaired electrons.
 According to the Pauli Exclusion Principle which states that no two
identical electrons may take up the same quantum state at the same time, the
electron spins are oriented in opposite directions.
 This causes the magnetic fields of the electrons to cancel out
 Thus there is no net magnetic moment,
and the atom cannot be attracted
into a magnetic field.
 In fact, diamagnetic substances are
weakly repelled by a magnetic field
as demonstrated with the pyrolytic carbon sheet

48. As an example,
Fe prefers to exist as Fe3+ and is known to have a coordination number of six.
Since the configuration of Fe3+ has five d electrons, we would expect to see five
unpaired spins in complexes with Fe.
This is true for [FeF6]3-; however, [Fe(CN)6]3−only has one unpaired electron,
making it a weaker magnet.
we expect CN−CN− ligands to have a stronger electric field than that
of F−F− ligands, so the energy differences in the d-orbitals should be greater for the
cyanide complex.

49. FINDING MAGNETIC MOMENT OF THE COMPLEX
GREATER NUMBER GREATER MAGNETIC
OF MOMENT
UNPAIRED ELECTRONS
Magnetic moment of thee substance
containing n unpaired electrons is
approximately equal to the given formula.
where, n = no of unpaired electrons
BM = Bohr Magneton (unit) eh/4πmc
µ = √n(n+2) B.M

50. EXAMPLE

51. LIMITATIONNS OF CFT
It considers only the metal ion d orbitals gives no consideration at all to other
metal orbitals such as s, px, py, pz orbitals
It is unable to account satisfactorily for the relative strength of ligands for eg, it
gives no explanation as to why H2O us a stronger ligand than OH- in the spectro
chemical series
According to this theory, the bond between the metal and ligands are purely
ionic
It gives no account of the partly covalent nature of the metal ligand bonds
The CFT cannot account for the π bonding in complexes.

52. VBT CFT
The bonding between the metal to
ligand sis considered to be covalent as
well as ionic
The bonding is considered to be purely
electrostatic between the positive metal
ion and the negative ions or negative
ends of the ligands
It does not take into account the
splitting of d-orbitals of the metal
It takes into account the splitting of d-
orbitals of the metals
It involves the concept of hybridisation
of orbital of the central metal ion before
complex formation
It does not involve the concept of
hybridisation
The colour of complexes are not clearly
explained
The colour of complexes are explained
in terms of electronic transitions
between the various d-orbitals of
different energies
The magnetic properties are not clearly
explained
The magnetic properties of complexes
depend upon the splitting of d-orbitals
of the central cation in different crystal

53. John teller distortion
JAHN TELLER DISTORTION

54.  The Jahn-Teller effect is a geometric distortion of a non-linear molecular
system that reduces its symmetry and energy.
This distortion is typically observed among octahedral complexes where the
two axial bonds can be shorter or longer than those of the equatorial bonds
It states that,
Any non-linear molecular system with degenerate electronic state
will be unstable
And thus distortion takes place in these non-linear system
Leads to system of lower symmetry and lower energy and
degeneracy of energy states will be removed
 This is most commonly observed with
transition metal octahedral complexes
however, it can be observed in
tetrahedral compounds as well
 The octahedral complex will either
elongate or compress the ligand bonds
as shown in Figure below

55. ELONGATION
 When an octahedral complex exhibits elongation, the axial bonds are
longer than the equatorial bonds
 It is called Z-out.
 Elongation Jahn-Teller distortions occur ,
when the degeneracy is broken by the stabilization (lowering in energy)
of the d orbitals with a z component, while the orbitals without
a z component are destabilized (higher in energy)
 This is due to the dxy ,dxy and dx2−y2 orbitals having greater overlap
with the ligand orbitals, resulting in the orbitals being higher in energy.
 Since the dx2−y2 orbital is antibonding, it is expected to increase in
energy due to elongation.
The dxy orbital is still nonbonding, but is destabilized due to the
interactions.
 Jahn-Teller elongations are well-documented for copper(II) octahedral
compounds.

56. Figure : Illustration of tetragonal distortion (elongation)
for an octahedral complex.
Figure : Structure of octahedral
copper(II) fluoride

57. COMPRESSION
 When an octahedral complex exhibits compression, the axial bonds are
shorter than the equatorial bonds
 Compression Jahn-Teller distortions occur
when the degeneracy is broken by the stabilization (lowering in energy)
of the d-orbitals without a z component, while the orbitals with
a z component are destabilized (higher in energy)
This is due to the z-component d orbitals having greater overlap with the
ligand orbitals, resulting in the orbitals being higher in energy.
 Since the dz
2 orbital is antibonding, it is expected to increase in energy due
to compression.
 The dxz and dyz orbitals are still nonbonding, but are destabilized due to the
interactions

58. Jahn-Teller distortion of octahedral Cu 2+. When the unpaired electron resides in the d z 2
orbital (left), the octahedra is compressed along the z direction. Alternatively, an unpaired
electron in the d x 2 −y 2 orbital (right) elongates the octahedra.
Source publicatio

59. SYMMETRICAL t2g and eg ORBITAL
 Vacant (t2g 0 , eg 0) , half filled (t2g3, eg2) or fully filled (t2g6 eg4)
configurations are termed as symmetrical t2g and eg orbitals
UNSYMMETRICAL t2g& eg ORBITALS
 The configurations other than vacant, half or fulfilled are termed as
unsymmetrical configuration
 eg., t2g1, t2g2, t2g4, t2g5, eg3, eg2 (low spin dx2-y2 ,dz2 0 OR dx2-y2 0 ,dz2 2)
Jahn teller distortion can be shown as
I. t2g (sym) – eg (sym) Regular octahedral ( NO JT distortion)
II. t2g(asym) – eg (sym) Slight distortion
III. t2g (sym) - eg (asym) very strong distortion

60. Larger ∆ low spin complexes
Figure : Electron configuration diagram of octahedral complexes (red indicates no
degeneracies possible, thus no Jahn-Teller effects).
The electron configurations highlighted in red (d3, low spin d6, d8, and d10) do
not exhibit Jahn-Teller distortions.
On the other hand d1, d2, low spin d4, low spin d5, low spin d7, and d9, would be
expected to exhibit Jahn-Teller distortion.
Some common examples include Cr3+, Co3+, and Ni2+.

61. SMALLER ∆ HIGH SPIN COMPLEX
Figure: High spin octahedral coordination diagram (red indicates no degeneracies
possible, thus no Jahn-Teller effect
The electron configurations highlighted in red (d3, high spin d5, d8, and d10) do
not exhibit Jahn-Teller distortions.
In general, degenerate electronic states occupying the eg orbital set tend to show
stronger Jahn-Teller effects.
This is primarily caused by the occupation of these high energy orbitals.

62. THANK YOU