Coordination chemistry - CFT

V.SANTHANAM
DEPARTMENT OF CHEMISTRY
SCSVMV

LIMITATIONS OF VBT
The valence bond approach could not explain
the following
 Electronic spectra
 Magnetic moments of most complexes.
So a more radical approach was put forward
which had only room for electrostatic forces

CRYSTAL-FIELD THEORY
 Model explaining bonding for transition metal
complexes
 Originally developed to explain properties for
crystalline material
 Electrostatic interaction between lone-pair
electrons result in coordination.

CFT assumptions
 Separate metal and ligand high energy
 Coordinated Metal - ligand stabilized
 Destabilization due to ligand -d electron
repulsion
 Splitting due to octahedral field.

Crystal Field Theory
 The electron pairs on the ligands are viewed
as point negative charges
 They interact with the d orbitals on the
central metal.
 The nature of the ligand and the tendency
toward covalent bonding is ignored.

Approach of ligands – Oh field

Crystal Field Theory
The repulsion
between ligand lone
pairs and the d
orbitals on the metal
results in a splitting of
the energy of the d
orbitals.

Crystal field theory
d-orbitals align along the octahedral axisd-orbitals align along the octahedral axis
will be affected the most.will be affected the most.
More directly the ligand attacks the metalMore directly the ligand attacks the metal
orbital, the higher the energy of the d-orbital, the higher the energy of the d-
orbital.orbital.
In an octahedral field the degeneracy of theIn an octahedral field the degeneracy of the
five d-orbitals is liftedfive d-orbitals is lifted

Ligand approach octahedral field – eg set

Ligand approach octahedral field – t2g
set

Splitting of the d-Orbitals
 The dThe dz2z2 and dand dx2-y2x2-y2 orbitals lie on the sameorbitals lie on the same
axes as negative charges.axes as negative charges.
 Therefore, there is a large, unfavorableTherefore, there is a large, unfavorable
interaction between ligand (-) orbitals.interaction between ligand (-) orbitals.
 These orbitals form the degenerateThese orbitals form the degenerate
high energy pair of energy levels.high energy pair of energy levels.

d Orbital Splitting
 In some texts and articles, the gap in the d
orbitals is assigned a value of 10Dq.
 The upper (eg) set goes up by 6Dq, and the
lower set (t2g) goes down by 4Dq.
 The actual size of the gap varies with the
metal and the ligands.

 The dThe dxyxy , d, dyxyx and dand dxzxz orbitals bisect theorbitals bisect the
negative charges.negative charges.
 Therefore, there is a smaller repulsionTherefore, there is a smaller repulsion
between ligand & metal for these orbitals.between ligand & metal for these orbitals.
 These orbitals form the degenerate lowThese orbitals form the degenerate low
energy set of energy levels.energy set of energy levels.

d-orbitals not pointing directly at axis are least
affected (stabilized) by electrostatic interaction
d-orbitals pointing directly at axis are
affected most by electrostatic interaction

d Orbital Splitting
________
Spherical
field
__ __
dz2 dx2-y2
__ __ __
dxy dxz dyz
∆o
+ 0.6∆o
- 0.4∆o
Octahedral field
eg
t2g
Free ion

Splitting pattern – Oh field
 The energy gap isThe energy gap is
referred to asreferred to as
∆ο (10 Dq)
 Also known asAlso known as
crystal field splittingcrystal field splitting
energyenergy

Factors affecting the magnitude of splitting
 Many experiments have shown that the magnitude
of splitting is depending upon both metal and
ligands.
JORGENSON’S RELATION
∆ο =∆ο = f . gf . g
f – metal parameter
g – ligand parameter

Metal factors
 Charge on the metal ion
 Number of d- electrons
 Principle quantum number of the metal d
electron

Number of d electrons - I
 Different charges same metal (No. of d electrons)
[Fe(H2O)6]2+
- 3 d 6
- 10,400 cm-1
[Fe(H2O)6]3+
- 3d5
- 13,700 cm-1
[Co(H2O)6]2+
- 3 d7
- 9,300 cm-1
[Co(H2O)6]3+
- 3d6
- 18,200 cm-1

Number of d electrons - II
 Same charge different metal ions
[Co(H2O)6]2+
- 3 d7
- 9,300 cm-1
[Ni(H2O)6]2+
- 3 d8
- 8,500 cm-1

Charge on the metal ion
 Same charge different metal
[V(H2O)6]2+
- 3d3
- 12,400 cm-1
[Cr(H2O)6]2+
- 3d3
- 17,400 cm-1

summary
For complexes having same geometry and
same ligands the crystal field splitting
 Increases with the increase in charge on the
ion (Same number of d - electrons)
 Decreases with increasing number of d -
electrons (Same charge on the ion)

Principle quantum number
 With increasing n value the splitting increases
[Co(NH3)6]3+
- 3d6
- 23,000 cm-1
[Rh(NH3)6]3+
- 4d6
- 34,000 cm-1
[Ir(NH3)6]3+
- 5d6
- 41,000 cm-1

Effect of ligand field
strength
Weak field Free ion strong field

Crystal field stabilisation
energy
 Already it is seen that t2g levels are lowered
while eg levels are raised in energy.
 The d – electron and ligand repulsion only
increases the energy.
 But the energy content of the system must
be a constant.
 So to maintain the centre of gravity the t2g
levels are getting lowered to an equivalent
amount.

+ 0.6 ∆o
- 0.4 ∆o
eg
t2g
Total energy change = 2 x (+ 0.6∆o) + 3 (- 0.4 ∆o) = 0

Crystal field stabilisation
energy
 Depending upon the field created by the
ligands the electrons are occupying the
various orbitals available.
 When t2g levels are getting filled the system is
getting lowered in energy
 Energy content increases if eg levels are filled
 If both of them are filled then the difference
between increases and decrease in energy is
calculated which is called crystal field
stabilisation energy

CFSE
 Gain in energy = + 0.6 ∆o x p
 Loss in energy = - 0.4 ∆o x q
 Net change in energy = [+ 0.6 x p + - 0.4 x q] ∆o
∆o = 10Dq
CFSE = [ -4Dq x q + 6Dq x p]

Splitting and Pairing
energy
 Pairing energy is the energy required for
accommodating second electron as a spin
pair to the first one in an orbital, against the
electrostatic repulsion.
 When the ligands are stronger, the splitting
of d orbitals is high.
 If splitting energy is more than the pairing
energy then according to Hund’s rule the
incoming electrons start to pair in the t2g
level itself

 .
 Fourth e- has choice:
 Higher orbital if ∆ is small; High spin
 Lower orbital if ∆ is large: Low spin.
 Weak field ligands - Small ∆ - High spin complex
 Strong field Ligands -Large ∆ - Low spin complex

d1
–d3
systems
Weak field Free ion strong field
∆ο
∆ο
CFSE
d1
- - 4Dq
d2
- - 8Dq
d3
- - 12Dq

Weak field d4
Free ion strong field
CFSE - -6Dq CFSE - - 16Dq + P

Weak field d5
-Free ion strong field
CFSE - - 20Dq + 2PCFSE - 0 Dq

Weak field d6
-Free ion strong field
CFSE - - 24Dq + 3PCFSE - -4Dq +P

Weak field d7
CFSE - - 18Dq + 3PCFSE - -8Dq +2P

Weak field d8
CFSE - - 12Dq + 3PCFSE - -12Dq +3P

Weak field d9
CFSE - -6Dq +4P

Weak field d10
CFSE - 0Dq +5P

Crystal Field Stabilization Energy
The first row transition metals in water are all
weak field, high spin cases.
dn
CFSE dn
CFSE
1 -4Dq 6 -4Dq + P
2 -8Dq 7 -8Dq + 2P
3 -12Dq 8 -12Dq + 3P
4 -6Dq 9 -6Dq + 4P
5 0 10 0 + 5P

High Spin vs. Low Spin
3d metals are generally high spin complexes
except with very strong ligands. CN-
forms
low spin complexes, especially with M3+
ions.
4d & 5d metals generally have a larger value of
∆o than for 3d metals. As a result, complexes
are typically low spin.

Colour of the complex
 The colors exhibited by most transition
metal complexes arises from the splitting of
the d orbitals.
 As electrons transition from the lower t2g set
to the eg set, light in the visible range is
absorbed.

Colour of the complexes
 The splitting due to the
nature of the ligand
can be observed and
measured using a
spectrophotometer
 Smaller values of ∆o
result in colors in the
green range. Larger
gaps shift the color to
yellow.

Spectrochemical / Fajan –Tsuchida
series
 Depending on the ligands present in a
complex the splitting value varies.
 By taking a particular metal, in a fixed
geometric field, the ligands are arranged in
the increasing order of the splitting caused
by them

Spectrochemical / Fajan –Tsuchida
series
I-
< Br-
<S2-
<Cl-
< NO3
-
< N3
-
< F-
< OH-
<
C2O4
2-
< H2O < NCS-
< CH3CN < pyridine <
NH3 < en < bipy < phen < NO2
-
< PPh3< CN-
< CO
Field Strength increases

Color Absorption of Co3+
Complexes
 The Colors of Some Complexes of the CoThe Colors of Some Complexes of the Co3+3+
IonIon
The complex with fluoride ion, [CoFThe complex with fluoride ion, [CoF66]]3+3+
, is high spin and has one absorption band., is high spin and has one absorption band.
The other complexes are low spin and have two absorption bands. In all but oneThe other complexes are low spin and have two absorption bands. In all but one
case, one of these absorptionsis in the visible region of the spectrum. Thecase, one of these absorptionsis in the visible region of the spectrum. The
wavelengths refer to the center of that absorption band.wavelengths refer to the center of that absorption band.
Complex IonComplex Ion Wavelength ofWavelength of Color of LightColor of Light Color of ComplexColor of Complex
light absorbedlight absorbed AbsorbedAbsorbed
[CoF[CoF66]] 3+3+
700 (nm)700 (nm) RedRed GreenGreen
[Co(C[Co(C22OO44))33]] 3+3+
600, 420600, 420 Yellow, violetYellow, violet Dark greenDark green
[Co(H[Co(H22O)O)66]] 3+3+
600, 400600, 400 Yellow, violetYellow, violet Blue-greenBlue-green
[Co(NH[Co(NH33))66]] 3+3+
475, 340475, 340 Blue, violetBlue, violet Yellow-orangeYellow-orange
[Co(en)[Co(en)33]] 3+3+
470, 340470, 340 Blue, ultravioletBlue, ultraviolet Yellow-orangeYellow-orange
[Co(CN)[Co(CN)66]] 3+3+
310310 UltravioletUltraviolet Pale YellowPale Yellow

The Spectrochemical Series
The complexes of cobalt (III) show the shift in color due to the ligand.
(a) CN–
(b) NO2
–
(c) phen (d) en (e) NH3
(f) gly (g) H2O (h) ox2–
(i) CO3
2–

Experimental Evidence for CFSE
 The hydration energies of the first row
transition metals should increase across the
period as the size of the metal ion gets smaller.
M2+
+ 6 H2O(l)  M(H2O)6
2+
 The heats of hydration show two “humps”
consistent with the expected LFSE for the metal
ions. The values for d5
and d10
are the same as
expected with a LFSE equal to 0.

Experimental Evidence for CFSE

Experimental Evidence of CFSE
ddoo
dd11
dd22
dd33
dd44
dd55
dd66
dd77
dd88
dd99
dd1010
LFSELFSE
In termsIn terms
ofof ΔΔoo
00 .4.4 .8.8 1.21.2 .6.6 00 .4.4 .8.8 1.21.2 .6.6 00

Structure of spinels
 Spinels are mixed oxides having general
formula AB2O4
 A is a divalent metal ion i.e. - A2+
 B is a trivalent metal ion i.e. - B3+
 The metals A and B may be same or different
 In spinels the oxide ions are arranged in cubic
close packed lattice

 In such situation each oxide ion will have 12
neighboring oxide ion at equidistant
 The lattice contains two types of coordination
sites
 Octahedral holes- surrounded by six oxide
ions – one hole per one oxide ion
 Tetrahedral holes –surrounded by four oxide
ions – two holes per one oxide ion

 Number of tetrahedral holes is twice the
number of octahedral holes.
 There are three types of spinels
 Normal
 Inverse
 Partially inverse

Normal spinel
 All the divalent cations occupy one of the eight
available tetrahedral holes
 Trivalent cations occupy the octahedral holes
 Represented as A2+
[B3+
2]O4
 Examples: FeCr2O4, Mn3O4, FeCr2S4, ZnAl2S4 and
ZnCr2Se4

Inverse spinels
 All divalent ions and half of the trivalent ions
occupy octahedral holes and other half of the
trivalent cations in the tetrahedral holes.
 Represented as B3+
[A2+
B2+
]O4
 Examples: CuFe2O4, MgFe2O4, Fe3O4, TiMn2O4,
TiFe2O4, TiZn2O4 and SnZn2O4

Partial inverse spinels
Examples of partially inverse spinel structures
include MgFe2O4, MnFe2O4 and NiAl2O4

Reason for inversion
 Let us consider the mixed oxides Mn3O4
(Normal spinel) and Fe3O4 (inverse spinel)
 Oxide ions are creating a weak field
 The table shows values of CFSE for the ions in
different sites
Site Mn3+
(d4
) Mn2+
(d5
) Fe3+
(d5
) Fe2+
(d6
)
Octahedral 6Dq 0 0 4Dq
Tetrahedral 1.78Dq 0 0 2.67Dq

 The values clearly shows that both trivalent
ions are having higher CFSE values at
octahedral holes.
 So preferably they tend to occupy the
octahedral sites
 This makes all the Mn3+
ions to occupy the
octahedral sites and Mn2+
ions in tetrahedral
sites
 Thus Mn3O4 is a normal spinel

 In the case of Fe3O4 , Fe3+
ions are expected to
be in the tetrahedral holes .
 Fe3+
ion in an octahedral hole is having higher
CFSE.
 So half of them are occupying octahedral
sites making the structure inverse

Stabilisation of oxidation state
 By using the CFSE values the stability of certain
oxidation state of a metal can be explained.
 In aqueous solutions Co2+
is stable and Co3+
is not
formed easily
 This is a direct consequence of higher CFSE for
[Co(H2O)6]2+
(d7
) [-8Dq]than [Co(H2O)6]3+
(d8
) [-4Dq]
 Similarly [Co(NH3)6]3+
(d6
) [-24Dq] has higher CFSE
than [Co(NH3)6]2+
(d7
) [-18Dq] so it is more stable.

Stereochemistry of
complexes
 Based on CFSE values we can say that why
Cu2+
is forming only square planar complexes
rather than octahedral
 SP symmetry complex has higher CFSE
Cu2+
in SP CFSE = 1.22 ∆o
Cu2+
in Oh CFSE = 0.18 ∆o

Coordination chemistry - CFT

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Coordination chemistry - CFT