This document provides an overview of bonding models for transition metal compounds. It discusses valence bond theory and how ligands donate electrons to empty metal orbitals to form bonds. Crystal field theory and ligand field theory are introduced to explain how the electronic structure of transition metals is affected by ligands in their coordination sphere. The ligands create a crystal field that splits the degenerate d-orbitals of the metal into different energy levels. The size of this splitting depends on factors like the metal ion and type of ligands. Spectroscopic properties of complexes, such as color, are determined by the energy gap between these split d-orbital levels.

1. Bonding Models
for
Transition metal compounds
Part 1
1
Christoph Sontag PhD, Phayao
University Aug. 2013
e-mail: c.sontag@web.de
Office SCI 2201

2. 2
 Review: electronic structure and periodic table
 VB Theory
 Crystal Field Theory
 Spectrochemical Series
 UV/VIS spectroscopy
 Ligand Field Theory

3. 3
Where are metals and non-metals in the periodic table ?
Where are the transition metals ?
Which elements form cations and which anions ?

4. How atoms form bonds ?
Watch tutorial videos on my website
http://myphayao.com  “Chemical Bonding”
Atoms use their VALENCE ELECTRONS to
exchange or share electrons to form bonds !
Examples: how many VE are in these elements ?
Na Ca Cl Ar C N S P
Which rule can we find for the main group
elements ?
4

5. Transition Metals have
valence electrons
in d-orbitals !!
5
How many groups of transition metals exist ?
(compare to the 8 MAIN GROUP elements)
And why ?

6. Number of Valence Electrons:
The 2 s electrons in each row will stay in d-orbitals
when we make a compound !!

7. Valence Bond (VB) Theory
Explain Metal-Ligand interaction:
• Each ligand gives 2 electrons to an empty
metal orbital and forms a covalent bond
• The metal orbitals have to be hybridized in
order to fit with the surrounding ligands
• We have to find out the right combination of
metal orbitals (all empty !) to combine with a
number of ligands (2, 4, 5 or 6)

8. Different idea from main group chemistry !
Here we assume that the central atom contributes 1 electron
and the ligands also 1 each
but in coordination chemistry, the electrons come all from the
ligands !

11. “Strong” ligands as CN(-) cause single electrons in d-orbitals
to combine together ! (forming “low-spin” compounds)

12. Similar to the Carbon sp3 in CH4

13. Important Hybrid orbitals

14. Exercises
What kind of complex will Ni(2+) form ?
(a)With weak ligands like Cl(-)
(b)With strong ligands like CN(-)

15. Possible answers:
(a) Consider [NiCl4]2+:
Ni0 = 4s2 3d8  Ni2+ = 4s0 3d8
3d 4s 4p
sp3

16. (b) Consider [Ni(CN)4]2-:
Ni0 = 4s2 3d8  Ni2+ = 4s0 3d8
3d 4s 4p
dsp2 4pz3d
Strong ligands
cause electron
pairing

17. Octahedral compounds
We need 6 empty hybrid orbitals on the metal
-> for octahedral shape, we need:
(a) d2sp3 (“inner shell”) or (b) sp3d2 (“outer shell”)
Example: [CoF6]3- (weak ligands)
How would [Co(NH3)6]3+ look like ?

18. Transition metal complexes
One interesting property of TM compounds is
their colour in solution:
Show Cu(II) experiments on:
http://www.youtube.com/watch?v=deNWxchzDRg

19. d-Orbitals
19
Review of orbitals

20. Energy Levels of d-Orbitals
“Crystal Field Splitting”
20
In an atom, the d-orbitals have the same energy (“degenerated”)
BUT: if molecules with high electron density approach the atom,
then the energy levels for these 2 d-orbitals go up:
el
el
el
el
el
el

21. 21
Look from the top down:

22. Energy Change in “crystal field”
22
el
el
el
el x2-y2

23. 23
“Crystal Field Splitting Theory”
Combination of a metal ion (positive)
and negative ligand molecules:
1. Electrostatic attraction
2. Metal-d electrons repulse ligand
electrons
3. Depending on the geometry of the
ligand sphere, metal d-orbitals are
split in different energy levels

24. 24

26. Colour of complexes
26
E = h * ν
Depends on the energy difference between the lower and higher
metal d-orbital levels !
Visible light is absorbed and pushes electrons up
=> The higher Δ, the more “blue” is the light absorbed

27. 27
Green colour => red is absorbed
Yellow colour => blue is absorbed
Conclusion:
The energy absorbed by the
green complex is lower than
by the yellow complex.
Calculate the absorbed light energy – example red light absorbed
(700 nm wavelength)
E = h * c / λ = 1240 eV * nm/ 700 nm = 1.77 eV
= 2.82 * 10^(-19) J
Blue light absorbed (400 nm wavelength) E =

28. Energy in electon volt eV
the amount of energy gained (or lost) by the charge of a
single electron moved across an electric potential difference of one volt.
Thus it is 1 volt (1 joule per coulomb, 1 J/C) multiplied by the negative of
the electron charge (−e, or −1.602176565(35)×10−19 C).
el
1 volt
Visible light photons
have an energy of
1.5 – 3.5 eV

29. Energy in spectroscopy
Energy of one photon: E = h * c / λ
Expressed as eV: E = 1240 eV nm / λ
Expressed in cm-1 E ~ 107 / λ
Expressed in J: E = 1.988×10−16 J nm / λ
Expressed in kJ/mol (for 1 mol of photons)
E = 1.988×10−19 x 6.022x1023 / λ =
= 1.2 x105 nm / λ
Calculate the energy of light of 500 nm in all 4 energies

30. Size of ∆
M
(+)
M
(n+)
low
high
L L
X(-) OH(-) H2O NH3 CN(-) CO

33. Estimate ∆
Co(3+) -> Ir(3+)
Fe(3+) -> Fe(2+)
Cr(F) 6
3- -> Cr(NH3)6
3+
Ni(NH3) 4
2+ -> Pd(NH3)4
2+

34. Spectrochemical Series
34
The splitting energy depends on:
1. metal ion:
high charge => high splitting : Ni(3+) > Ni(2+)
high period => high splitting : Pd > Ni
Pt4+ > Ir3+ > Rh3+ > Co3+ > Cr3+ > Fe3+ > Fe2+ > Co2+ > Ni2+ > Mn2+
2. ligands:
empirical order by measurements = “spectrochemical series”

35. Energy Calculations
UV/VIS spectra: scala in cm-1 = 1/λ = ν “wavenumber”
BECAUSE: the wavenumber ~ energy (h* ν)
10.000 cm-1 ≈ 120 kJ/mol => wavelength λ = ..............
Example: Fe(H2O)6
2+ = d ?
Compare the energy of low and high spin !
(Δ = 10.400 cm-1 / P = 17.600 cm-1)
Is the complex dia- or paramagnetic ?
35

36. 36
End of Part 1
Summary:
• Transition metal ions have d-electrons which can interact with electron
rich small molecules like ammonia, water, chloride ... (we call them
“ligands”)
• The electronic field of the ligands around the d-orbitals cause a splitting
in their energy -> for ML6 the d-orbitals split into 3 low- and 2 high-energy
orbitals
• The amount of splitting increases with higher charge on the metal and
with higher electron density on the ligands
• Electrons in the lower d-orbitals can jump to the high energy levels by
absorbing light energy => the complex has a colour !
THANKS FOR WATCHING … SEE YOU LATER !