[niche]

I obtained a UV/Vis spectrum (Abs vs. wavelength) for K3Fe(CN)6 dissolved in H2O.  I have distinct peaks at around 410, 299, and 259 nm.  The readings beyond the 259 wavelength skyrocketed off the charts, signaling an intense absorption in the UV region.  I believe it is due to the charge transfer bands, but I'm searching through references to help explain the exponential gain.  Also the 3 peaks listed above are characteristic to C=N ? Any help linking these peaks to the molecular structure would be greatly appreciated.

[Alpha-Omega]

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:

MnO₄^- : 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

CrO₄^2- : 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 MnO₄^-.

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., MnO₄^- where Mn is in the +7 oxidation state and CrO₄^2- 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(H₂0₄)]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(H₂O)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*  NO₃^-(313nm), CO₃^-2(217nm), NO₂^-(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 _x2-y2 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^- <C₂O₄^2- ~H₂O)<SCN^- <NH₃ <ethylenediamine<o-phenanthroline<NO₂^-<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
I₃^- (complex of I^- and I₂)

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 Fe³⁺ SCN^------>  Fe²⁺SCN⁰.  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

Solvents

Solvent must be transparent in wavelengths  being studied also may interact with chromophores. Polar solvents water, alcohols esters and ketones tend to obliterate fine structure.

In order to obtain empirical data for specific absorbances for CN^- you will have to consult UV-Vis library spectra.

[enahs]

It is just as likely that at that point it switched to a different source of light or grating, and was not calibrated properly and so it did not scale the background (or even a typical signal) correct.

[niche]

Thanks Alpha-Omega for the detailed response.

Also I found an explanation for the intense absorption in the UV region concerning the photo chemistry of complex ions.  It's stated by Mulliken that the extreme intensity of absorption in the UV area is due to the absence of hindering selection rules and to the large change in polarity accompanying the transition from the ground state to the excited state. 

He also states that the peaks within the visible range (d-d bands) have a much lower intensity because the transitions are parity forbidden and sometimes spin forbidden as well.

The spectrum is beginning to make more sense to me.  Can I get an Expert's opinion again on these selection rules?

[Alpha-Omega]

I am attaching 4 Word documents that have a pretty thorough explanation of selection rules.  I have put 1 and 2 in this post.  I had to attach 3 and 4 in separate posts due to size limitations/restrictions. Three posts total...

[Alpha-Omega]

Word Document 3 attached...

[Alpha-Omega]

Word document 4 attached...