V.SANTHANAM - Kar
Transcript of V.SANTHANAM - Kar
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V.SANTHANAMDEPARTMENT OF CHEMISTRY
SCSVMV
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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
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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.
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CFT assumptions
Separate metal and ligand high energy
Coordinated Metal - ligand stabilized
Destabilization due to ligand -d electron repulsion
Splitting due to octahedral field.
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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.
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d Orbitals
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Approach of ligands – Oh field
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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.
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Crystal field theory
d-orbitals align along the octahedral axis d-orbitals align along the octahedral axis will be affected the most.will be affected the most.
More directly the ligand attacks the metal More 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 the In an octahedral field the degeneracy of the five d-orbitals is liftedfive d-orbitals is lifted
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i
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Ligand approach octahedral field – eg set
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Ligand approach octahedral field – t2g set
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Splitting of the d-Orbitals
The dThe dz2z2 and d and dx2-y2 x2-y2 orbitals lie on the same orbitals lie on the same axes as negative charges.axes as negative charges.
Therefore, there is a large, unfavorable Therefore, there is a large, unfavorable interaction between ligand (-) orbitals.interaction between ligand (-) orbitals.
These orbitals form the degenerate These orbitals form the degenerate high energy pair of energy levels.high energy pair of energy levels.
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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.
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The dThe dxyxy , d , dyxyx and d and dxzxz orbitals bisect the orbitals bisect the negative charges.negative charges.
Therefore, there is a smaller repulsion Therefore, there is a smaller repulsion between ligand & metal for these orbitals.between ligand & metal for these orbitals.
These orbitals form the degenerate low These orbitals form the degenerate low energy set of energy levels.energy set of energy levels.
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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
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d Orbital Splitting
________
Spherical
field
__ __dz2 dx2-y2
__ __ __dxy dxz dyz
∆o + 0.6∆o
- 0.4∆o
Octahedral field
eg
t2g
Free ion
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Splitting pattern – Oh field
The energy gap is The energy gap is referred to as referred to as ∆ο (10 Dq)
Also known as Also known as crystal field splitting crystal field splitting energyenergy
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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 parameterg – ligand parameter
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Metal factors
Charge on the metal ion
Number of d- electrons
Principle quantum number of the metal d electron
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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
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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
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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
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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)
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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
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Effect of ligand field strength
Weak field Free ion strong field
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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.
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+ 0.6 ∆o
- 0.4 ∆o
eg
t2g
Total energy change = 2 x (+ 0.6∆o) + 3 (- 0.4 ∆o) = 0
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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
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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]
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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
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.
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
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d1 –d3 systems
Weak field Free ion strong field
∆ο∆ο
CFSE d1 - - 4Dqd2 - - 8Dq
d3 - - 12Dq
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Weak field d4 Free ion strong field
CFSE - -6Dq CFSE - - 16Dq + P
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Weak field d5 -Free ion strong field
CFSE - - 20Dq + 2PCFSE - 0 Dq
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Weak field d6 -Free ion strong field
CFSE - - 24Dq + 3PCFSE - -4Dq +P
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Weak field d7 Free ion strong field
CFSE - - 18Dq + 3PCFSE - -8Dq +2P
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Weak field d8 Free ion strong field
CFSE - - 12Dq + 3PCFSE - -12Dq +3P
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Weak field d9 Free ion strong field
CFSE - -6Dq +4P
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Weak field d10 Free ion strong field
CFSE - 0Dq +5P
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Crystal Field Stabilization Energy
The first row transition metals in water are all weak field, high spin cases.
dn CFSE dn CFSE1 -4Dq 6 -4Dq + P
2 -8Dq 7 -8Dq + 2P
3 -12Dq 8 -12Dq + 3P
4 -6Dq 9 -6Dq + 4P
5 0 10 0 + 5P
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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.
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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.
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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.
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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
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Spectrochemical / Fajan –Tsuchida series
I- < Br- <S2- <Cl- < NO3- < N3
- < F-< OH- <
C2O42- < H2O < NCS- < CH3CN < pyridine <
NH3 < en < bipy < phen < NO2- < PPh3 < CN- < CO
Field Strength increases
Field Strength increases
Field Strength increases
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Color Absorption of Co3+ Complexes The Colors of Some Complexes of the CoThe Colors of Some Complexes of the Co 3+ 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 one The 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. The case, 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 of Wavelength of Color of Light Color of Light Color of ComplexColor of Complex light absorbed light absorbed Absorbed Absorbed
[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, ultraviolet Blue, ultraviolet Yellow-orangeYellow-orange
[Co(CN)[Co(CN)66] ] 3+3+ 310310 Ultraviolet Ultraviolet Pale YellowPale Yellow
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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–
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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)62+
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.
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Experimental Evidence for CFSE
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Experimental Evidence of CFSEddoo dd11 dd22 dd33 dd44 dd55 dd66 dd77 dd88 dd99 dd1010
LFSELFSE
In terms In terms of of ΔΔoo
00 .4.4 .8.8 1.21.2 .6.6 00 .4.4 .8.8 1.21.2 .6.6 00
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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
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Structure of spinels
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
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Structure of spinels
Number of tetrahedral holes is twice the number of octahedral holes.
There are three types of spinels
Normal Inverse Partially inverse
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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
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Structure of spinels
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Structure of spinel - MgAl2O4
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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
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Partial inverse spinels
Examples of partially inverse spinel structures include MgFe2O4, MnFe2O4 and NiAl2O4
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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
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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
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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
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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.
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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 CFSECu2+ in SP CFSE = 1.22 ∆oCu2+ in Oh CFSE = 0.18 ∆o
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