Download F. ELECTRONIC SPECTRA OF COORDINATION COMPOUNDS

F. ELECTRONIC SPECTRA OF COORDINATION COMPOUNDS
Introduction:
Transition metal compounds display a wide variety of colors and many of these compounds are used
as artist's pigments (e.g., copper phthalocyanine, prussian blue, and Thénard's blue are all various
shades of blue). A compound is colored if it absorbs part of the visible light. When a sample absorbs
visible light, the color we observe is the sum of the remaining colors that are reflected or transmitted
by the sample and strikes our eyes. Remaining colors can conveniently be remembered as the color
pairs on opposite sides of the color wheel (Newton Wheel) shown in Figure. 1.
Figure.1 Newton Wheel (It was first created by
Sir Isaac Newton in 1666.)
The color of these compounds is due to electronic transitions induced by the incident light, causing
absorption at certain wavelengths, depending on the structure and bonding in the molecule, which is
also known as Crystal Field Theory (CFT) in Figure 2. UV–Visible spectroscopy is used to evaluate
the absorption properties of molecules.
The electronic transitions of transition metal compounds can be categorized into three types: charge
transfer, d–d spin–allowed, and d–d spin–forbidden. Charge transfer bands occur when the excited
electron moves from a metal centered orbital to a ligand–centered orbital (metal–to–ligand charge
transfer – MLCT) or when the excited electron moves from a ligand–centered orbital to a metal
centered orbital (ligand–to–metal charge transfer – LMCT). MLCT bands are more common than
LMCT. Spin–allowed and spin–forbidden d–d transitions find the excited electron moving from one d
orbital to another d orbital. In the spin–allowed transition the spin quantum number (or more
accurately, the spin angular momentum) in the ground and excited states are the same while in the
spin–forbidden transition the spin state changes during the excitation. These three types of transitions
are distinguished by the molar absorption coefficient (in the Beer–Lambert law: A=bc): charge
transfer bands have large , generally in the range of 1000 to 10000 L/mol–cm; spin–allowed d–d
transition have values of 10–100 L/mol–cm for octahedral complexes and 500–1000 for tetrahedral
complexes; and spin–forbidden bands have values less than 5 L/mol–cm. Spin–forbidden bands are
often not observed. UV–Vis spectroscopy can be used to distinguish octahedral from tetrahedral
geometries (Figure. 3). It is important to report both peak positions and molar extinction coefficients
for electronic spectra. Crystal Field Splitting of d- orbitals is arranged differently by changing the
environment around the central metal, which can be seen in Figure. 3.
Figure 2. Summary of Crystal Field Theory
The primary use of UV–Vis spectra for transition metal complexes is the elucidation of electronic
structure parameters. Generally this means the crystal field parameter ∆0 and sometimes the Racah
inter-electron repulsion constant B. These values are easy to find for most of the electron
configurations of octahedral complexes. For d1, d4 (high spin), d6 (high spin), and d9 complexes a
single spin–allowed d–d transition is observed with energy ∆0. High spin d5 complexes have no spin–
allowed d–d bands so are lightly colored unless charge transfer bands dominate the spectrum. For d 3
and d8 octahedral complexes, three spin–allowed transitions are predicted at energies given by:
v1 ∆0
v2 7.5B 1.5∆0 1/2 [225B2 10∆02 18B∆0]
v3 7.5B 1.5∆0 1/2 [225B2 10∆02 18B∆0]
where all the parameters are in units of cm–1. If all three transitions are observed, ∆0 can easily be
found from ν1 and B can be found from 15B=ν3+ν2–3ν1. Often a charge transfer band will obscure 3
so that B must be found from

v2 2v1v2 v1
3B 

5v2 9v1
If the complex has a d2 or d7 (high spin) configuration, three bands are also expected and are given by:
v1 7.5B 0.5∆0 1/2 [225B2 10∆02 18B∆0]
v2 7.5B 1.5∆0 1/2 [225B2 10∆02 18B∆0]
v3  [225B2 10∆02 18B∆0]
where again, all parameters are in units of cm–1. Here ∆0 is found from the difference of the first two
bands, ∆0=ν2–ν1, and B is found from 15B=ν3+ν2–3ν1. If only the first two transitions can be
observed, then B can be found from
v2 2v1 v1
3B 

9v1 4v2
For complexes with other electron configurations no simple formulas exist to find the crystal field
parameters.
The crystal field strength, ∆0, measures the splitting between the t2g and eg orbitals and depends on
both the metal ion and the ligand. The spectrochemical series denotes the variation of ∆0 as a function
of ligand with weak ligands such as Cl– or H2O leading to small values of ∆0 and stronger ligands like
NH3 or CN– giving larger values of ∆0. The inter-electron repulsion constant B is usually reported as
β= Bcomplex/Bfree ion. Bonding reduces the inter-electron repulsion in a complex relative to the free (gas
phase) ion (this happens because the metal d electrons expand to bond to the ligands and electrons that
are farther apart from each other repel each other less) so that βis always less than 1. The value of
βalso varies from ligand to ligand and this series is known as the nephelauxetic series, which is
similar to but not identical to the spectrochemical series.
Figure 3. Crystal Field Splitting parameters for different symmetrical fields.
The energy difference between spherical field and octahedral field is called Crystal Field Stabilization
Energy (CFSE), which can be calculated as:
CFSE = 0.4 ∆0 . (# of electrons in t2g level) - 0.6 ∆0 . (# of electrons in t2g level)
Also, CFSE would be different for high spin and/or low spin complexes.
Experimental reports of UV–Vis data are not complete unless both the peak position and the molar
absorptivity of each peak are given. A solution should be prepared so that the absorbance is between
1.5 and 2.0 and analytically diluted 4 or 5 times so that the most dilute concentration is at least 10
times smaller than the initial concentration. The spectrum is recorded at each dilution and the
absorbance of each peak is plotted as a function of concentration. This plot should yield a straight line
with a slope of (for a 1 cm cell) and an intercept of 0. If the slope has curvature, then the either
experimental technique needs to be improved or the compound is reacting with solvent, itself, or
perhaps air. To calculate ∆0/10 and/or B, the peak positions must be identified as spin–allowed d–d
transitions in an octahedral environment and then converted to energy units, usually cm–1.
Reference James E. Huheey, Inorganic Chemistry, Harper Collins College Publisher, 1993, p. 387.
Experiment 9.
Experimental Procedure:
1. Prepare the diluted solutions of
i) KMnO4
ii) [Co(NH3)5NO2]SO4 (product material of experiment #4 part-III),
iii) [Ni(NH3)6]Cl2 (product material of experiment #3 part-I),
iv) FeCl2.4H2O,
v) VO(acac)2 (product material of experiment #7),
vi) K[Cr(ox)2(H2O)2].2H2O
by using 0.1-0.2 g of samples in 50 mL water. If the concentration of the solution is high (dark color),
you need to dilute the solution.
2. Scan the UV-Vis range of 300 nm to 900 nm by using UV-Vis Spectrophotometer. The instruction
for this instrument is given by the lab intsructors. DO NOT USE THE INSTRUMENT BY
YOURSELF.
Calculating Crystal Field Splitting Parameter, ∆0:
1. Convert all transition wavelenght values (nm) to energy (kj/mol) by using E=hc/ and multiply by
Avagadro’s number to molar energy.
2. Calculate Racah Parameter by using equations given in the introduction section.
3. Calculate Crystal Splitting Parameter, ∆0, by using equations given in the introduction section
4. Calculate Crystal Field Stabilization Energy (CFSE) by using equations given in the introduction
section.
DATA and RESULTS SHEET
Sample
i) KMnO4
ii) [Co(NH3)5NO2]SO4
iii) [Ni(NH3)6]Cl2
iv) FeCl2.4H2O
v) VO(acac)2
vi) K[Cr(ox)2(H2O)2].2H2O
1
2
3
B
∆0
CFSE