I posted the answer to a question like this a few days ago. I will give you the same information. The objective is to give you enough information so you can make the determination yourself. Plus there is additional information in here I beleive you requested.
Additionally, I will also post some information regarding the Jahn-Teller Effect per your request....
Titanium has a number of oxidation states Ti
-1, Ti
o, Ti
+2, Ti
+3, and Ti
+4[Ti(H2O)6]2+
Is a d2 complex (Ti2+ = 1s2 2s2 2p6 3s2 3p6 3d2) with weak field ligands (High Spin/LFSE -4/5).
[Ti(CN)6]4-
Is a d2 complex (Ti2+ = 1s2 2s2 2p6 3s2 3p6 3d2) with strong field ligands (Low Spin/LFSE -4/5).Strong-field ligands = low-spin complexes. Strong field ligands have pi-acceptor orbitals or low-lying d-orbitals: p* as in CO or CN
–, p* as in CH
2=CH
2, low lying d as in PR
3, PF
3Weak field ligands = high-spin complexes. Weak field ligands have pi-donor orbitals: Usually multiple p-orbitals as in X–.
Intermediate field ligands = usually high-spin for +2 ions, low-spin for +3 ions. Intermediate field ligands have few, or no pi-donor or acceptor orbitals, or there is a poor match in energy of available pi-orbitals: NH
3, H
2O, pyridine.
A close relationship exists between the color of a substance and its electronic structure. A molecule or ion will exhibit absorption in the visible or ultraviolet region when radiation causes an electronic transition within its structure. Thus, the absorption of light by a sample in the ultraviolet or visible region is accompanied by a change in the electronic state of the molecules in the sample.
Transition metal complexes are highly colored. The colors exhibited by these complexes are due to electronic transitions by the absorption of light.
Most transitions related to metal complexes are either "d-d transitions" or "charge transfer bands." In a d-d transition, an electron in a d orbital on the metal is excited by a photon to another d orbital of higher energy.
A charge transfer band facilitates the promotion of an electron from a metal-based orbital into an empty ligand based orbital-MLCT (Metal to Ligand Charge Transfer).
Charge Transfer Transitions:
Some transition metal complexes with no d-electrons are colored. This is because there can be electronic transitions in the visible region that do not involve d-electrons:
MnO
4- : In this case, electrons in filled oxygen based orbitals are excited into empty d-orbitals. This type of Ligand to Metal Charge Transfer band gives rise to the intense purple color of permanganate
CrO
42- : The intense yellow color observed is due to the LMCT band.
Metal to ligand charge transfer bands also occur in the visible regions for some complexes. Charge transfer transitions are often much more probable than d-d transitions. Hence the intense color of MnO
4-.
MLCT complexes experience a partial transfer of electrons from the metal to the ligand. Usually this occurs for metals with highly-filled d orbitals capable of donating electrons into the antibonding orbitals of the ligand. The non-bonding d orbitals must match the antibonding orbitals in terms of size, shape, and symmetry
Conversely, there can be excitation of an electron in a ligand-based orbital into an empty metal-based orbital LMCT (Ligand to Metal Charge Transfer). These transitions can be observed/monitored by electronic spectroscopy such as UV-Vis.
LMCT complexes experience a partial transfer of electrons from the ligand to the metal. This is common when the metal has a high oxidation state (e.g., MnO
4- where Mn is in the +7 oxidation state and CrO
42- where Cr is in the +6 oxidation state).
The energy supplied by the light will promote electrons from their ground state orbitals to higher energy, excited state orbitals or antibonding orbitals.
CN
- is a pi acceptor (overlap of d, p*, and p-orbitals with metal d orbitals). The overlap is good with ligand d and p pi*-orbitals, poorer with ligand p-orbitals
The wavelength of maximum absorption and the intensity of absorption are determined by molecular structure. Transitions to π* antibonding orbitals which occur in the ultraviolet region for a particular molecule may take place in the visible region if the molecular structure is modified.
Many inorganic compounds in solution also show absorption in the visible region. These include salts of elements with incomplete inner electron shells (mainly transition metals) whose ions are complexed by hydration e.g. [Cu(H
204)]2+. Such absorptions arise from a charge transfer process, where electrons are moved from one part of the system to another by the energy provided by the visible light.
The ability to complex many metals, particularly the transition elements, with complex organic and inorganic molecules which absorb in the visible region provides the basis for their quantitative spectrometric analysis. The absorptions are due to movement of electrons between energy levels of the organo-metal complex. These complexing systems are termed spectrometric reagents (e.g., dithizone, azo reagents, PAN, thoron, zincon dithiocarbamate, 8-hydroxyquinoline, formaldoxime and thiocyanate.).
In addition, many inorganic ions in solution also absorb in the visible region e.g. salts of Ni, Co, Cu, V etc. and particularly elements with incomplete inner electron shells whose ions are complexed by hydration e.g. (Cu(H
2O)
4)
2+. Such absorptions arise from a charge transfer process where electrons are moved from one part of the system to another due to the energy provided by the visible light.
n -----> pi* and pi -----> pi* Transitions
Most UV absorption of organics due to these two transitions 200-700 nm. Both arise only in molecules with double bonds for B electrons
Characteristics
n -----> pi*
low absorptivities (10-100 L cm-1 mol-1). Shift to higher E (shorter wavelengths) in more polar solvents- called a hypsochromic or blue shift n electrons are solvated so are at lower E in polar solvent. Shift can be as much as 30 nm in water where have H bonding E of shift seems to be about the same as the E of the H bond. Upper end of transition unaffected by solvent.
Absorption by Inorganic Anions
A number of inorganic anions have n -----> pi* NO
3-(313nm), CO
3-2(217nm), NO
2-(360&280nm), CN
-(230nm), and trithiocarbonate(500nm).
pi -----> pi*
high absorptivities (1000-10,000). Often (but not always) shift to lower E (larger wavelength) in polar solvents- called a bathochromic or red shift. Much smaller than hypsochromics shift, only 5 nm. Both pi and pi* electrons slightly solvated and favored pi* is slightly more strongly favored so overall transition E decreased as polarity increases.
1st and 2nd transition metal series tend to absorb UV radiation in attained oxidation state. The bands are broad, strongly influenced by ligands. Transition involves moving electron between d orbitals.Two theories used to rationalize observed colors: Crystal-Field Theory (simpler) and Molecular Orbital Theory (more complex, but better numbers).
Theory
5 d orbitals in the absence of external magnetic or electric field are degenerate (all have the same E). The electrons can move freely between orbitals. In a complex (or in solution) ligands molecules have electrons, hence there is an external magnetic field. In an external field, some d’s have different E and can absorb E as electrons move.
5 d orbitals: d
xy, d
xz and d
yz similar shape, and are oriented between axes. The d
x2-y2 and d
z2 orbitals are oriented along the axes. The most common complex octahedral, Where one ligand is oriented along each axis. This places electrons and electric field at axes. This results in a minimal, equal, effect on the d
xy, d
xz and d
yz orbitals and will raise the E of the d sub]x2-y2 [/sub] and d
z2 orbitals.
This change in E results in a difference in E of delta. The magnitude of delta depends on several factors:
charge of metal, position of atom in periodic table, ligand field strength (property of ligand), high ligand field strength – large delta- shorter wavelength.
Ligand Field Strengths
I
- <Br
- <Cl
- <F
- <OH
- <C
2O
42- ~H
2O)<SCN
- <NH
3 <ethylenediamine<o-phenanthroline<NO
2-<CN
-Charge Transfer Absorption
Very strong absorptions, so sensitive means to detect and quantitate. Many inorganic complexes exhibit charge transfer absorption; and, are called charge transfer complexes.
Examples:
Fe
+3 SCN
- complex
Fe
+3 Phenol complexes
Fe
+2 0
-penanthroline
I
3- (complex of I
- and I
2)
One group must be strong electron donor (Lewis base); and, the other must be strong electron acceptor (Lewis acid). Transition involves moving electron from donor to acceptor. Think of it as internal redox reaction.
Note: Other transition have studied, moved electron between molecular orbitals, did not move from one atom to another Fe
3+ SCN------> Fe
2+SCN
0. Usually electron returns to it’s original state quickly. If the complex dissociates before this happens, then you have a photochemical redox reaction. Most of the time the metal is the electron acceptor (but not always). Some organic molecules can form charge transfer complexes,
I am attaching 4 word documents that outline the selection rules and how they affect UV-Vis spectra for coordination complexes. Due to sise restrictions I will have to do that in 3 additional posts.
I have also attached the information regarding the Jahn Teller Theorem/Effect you requested.
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Topic: Absorption Spectra (Read 18620 times)