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MOLECULAR ORBITAL THEORY (MOT) - Coggle Diagram
MOLECULAR ORBITAL THEORY (MOT)
Classification of overlapping
σ- overlapping ( σ - bond)
Two orbitals are directly overlapped each other (crossed axis).
The orbital must be exist on every bonding and it refer as an orbital to produce skeleton bond between two bonded atoms.
Lead to the formation of σ-bonding when overlapped are
- s-s
- p-s
- Px-Px
π-overlapping ( π - bond)
Overlapping forms when two orbitals are approached together through
side position
.
Not necessarily present in each bonding
between two bonded atoms but depends on other position orbital of each bonded atoms after s-orbital done.
The presence of the orbital enhancing the strength of bond in the two bonded atoms.
Lead to the formation of π - orbital when overlapped are
- Py-Py
- Pz-Pz
- d-p
μ- overlapping (μ - bond)
Diagonally overlapping
sp^3-dxy
p-d
Always found in metal-bridging ligand in complex.
δ- overlapping (δ-bond)
dxy-dxy
Metal-metal bond
Always found in cluster complex.
To support disadvantages of Valence Bond Theory (VBT) & Crystal Field Theory (CFT)
The theory has been introduced
The overlapping concept between orbitals of metal and ligand similar to that in the other molecules.
Overlapping must be occur between two components which is the metal and its orbital :
an equal / nearly equal energy levels.
a position towards overlapping to each other.
The number of new orbitals produced = The number of orbital used in overlapping.
For example : 6 orbitals involved in overlapping = 6 new orbitals
Classified in 4 groups depends on the position of overlapping done
σ (sigma) bonding
π (pi) bonding
μ (miu) bonding - bridging ligand
δ (delta) bonding
ML6 complex : σ - overlapping only
Other 3 orbitals of metal (dxy, dxz, and dyz) are not in same direction and do not undergo overlapping and geometrically unchanged.
3 orbitals (n+1)p, 1 orbital (n+1)s of metal have almost an equal energy with ligand orbital.
Example of these ligands : ammonia, NH3 and en (ethylenediammine)
6 lone pair orbitals of ligand and 2 orbitals of metal (dx^2-y^2 and dz^2) are in same direction and having almost equal energy.
nd, (n+1)s and (n+1)p orbitals on the metal overlap with 1 orbital on each of the six ligands.
Forms 15 molecular orbital
Six are bonding, energies are lower.
Six are anti-bonding, energies are higher.
Three are non-bonding having same energy.
In a complex
ML6
which only undergo
σ-overlapping only
,
the total number of electrons = total of number of electrons of d-orbital of metal + 12 electron of ligands which is 6 orbitals involved.
ML6 complex : σ and π-overlapping
The side overlapped will results in three orbitals dxy, dyz and dxz not a "non-bonding" character.
The σ-overlapping is dx^2-y^2 and dz^2.
Form σ-overlapping and also perform π-overlapping (side position) withy d-orbitals (dxy, dyz and dxz)
Two types of π-overlapped which are π-donor and π-acceptor.
-The existence of the π-overlapped will have an effect on the stability of the complex.
π-Acceptor Ligand
An empty orbital of ligand (have orbital but does not contain any electron)
Has higher energy level than dxy, dyz, and dxz (t2g).
- Example : CN- , CO, NO2 and PPh3
contains *π orbitals which can overlap with the dxy, dyz and dxz of the metals by π bond
There is no electron in the empty
π orbital,
*the total number of electrons = the electrons in the d-orbital of the metal + 12 electrons from the ligand.
Example : [Co(CN)6] ^ (-3)
Δ0 of t2g and *eg increases.
Form
low spin
.
More stable electron fills the t2g orbital & has a lower energy.
There is a bonding of electron from metal ---> ligand called backbonding in the π-bonding.
π-Donor Ligand
Having two or more lone pairs of electrons with one of these lone pair having σ-bonding while others can perform the π-bonding with dxy, dxz and dyz orbitals.
The lone pair of electron has a lower energy than dxy, dyz and dxz (t2g) orbitals.
The direction is suitable for overlapping to occur.
- Examples: F-, OH-, Cl- Br- (weak ligands)
The ligand supply electrons in the π-overlapped and electrons flow form ligand ---> metal
The total number of ligand electrons involved in bonding is
higher depends on the total π-electrons.
Example : The total number of electrons donating by ligands in [Co(F)6] ^ (-3) complex is 18 ( 6 σ and 3 π ) and d-orbitals from metal.
Total number of electron = 6 + 18 = 24
Δ0 of *t2g
and *eg increases.
Form
high spin
(unpaired)
Unstable since the d electrons filled in anti-bonding *t2g & has higher energy.
ML4 : Tetrahedron Complexes
Produces
very small Δt compared to Δ0.
Δt is the energy difference of *t2g and eg.
Have high spin.
Complex yield is
unstable.
- Tetrahedral complex yield with ML4 with d^1 - d^7 , d^9 and d^10 with all ligands.
d^8 complex
except Pt(II) forms tetrahedral complex with weak ligands only.
Example : [CoF4]- , [NiCl4] ^ (-2)
ML4 : Square Planar Complexes
Obtained with
d^8 complex with certain ligands only.
Pt(II) forms square planar complex with all ligands. (Larger size of ion)
Ni(II) forms square planar complex with strong ligands only.
For example : CN- , CO, NO2 , PPh3 (Posses high splitting energy).
More stable low spin complex is formed.
Advantages Of MOT
Able to explain the electronic configuration of the complexes.
Able to explain the strength of the electro spectrochemical series.
Able to explain the color yield from the complex which due to electronic transition of t2g and eg.
(d-transition)
.
Able to explain the stability pf complexes with respect to the size of Δ0. The larger the value of Δ0, the more stable the complex.
Measurement Of The Effect of π Bonding
Carbonyl complex
is chosen due to
---> C ≡ O group has strong IR absorption and easily identified (due to triple bond character of CO)
---> Absorption frequency range is away from other organic functional groups which make it easier to identify.
Absorption frequency is proportional with the bonding strength in C ≡ O.
The stronger the bond give higher frequency.
The existence of M-CO π-bond, effect the C-O bond strength.
π Back Bonding
No Back
3CO ligands share the π electron density makes M-C bonds become stronger & the C ≡ O become weaker.
M ---------> L
If the ligand can
accept electron density from metal, for example : PPh3 , CN-
The
electron density distributed equally among these 6 ligands -----> no change in absorption frequency.
Charge effect on the central atom
The more (+) metal ion, less tendency to donate electron for back bonding to *π orbital.
The M-C bond become weaker
and
C≡O bond become stronger, the absorption frequency is higher.
Effect Of The Metal Charge