Hi Tarunpreet, I think n-pi* do occur in carbonyls but the energy of that transition is higher so does not fall in the visible. As usual with spectroscopy there is the allowedness of the transition and its energy. The colour comes from the energy gap. If this is a discrete molecular phenomenon then it should not change with concentration (exept to increase in intensity due ot Beer-Lambert law) but if it is due to aggregation then new bands should be observed.

Thiocarbonyl compounds  absorbed in visible (n->pi) see publication, not carbonyl compounds

I disagree with Tarunpreet; general consensus as a main group chemist is that sulfur does not implement d-orbitals in bonding (but can use symmetry equivalent linear combinations of sigma*).

The likely transition giving rise to colour is an n --> pi* transition. In terms of MO theory the poorer pi bonding between C and S leads to a lower pi* (than C=O). Conversely the S lone pair (3p) is higher in energy than the O (2p). Hence the n->pi* transition is lower in energy in C=S systems than C=O, appearing towards the visible rather than UV.  

Many thanks! I have to admit we also considered the polarizability argument put forward by Tarunpreet, but our X-ray and comp geometry optimization data did not support any increased CT contribution.  Didier, many thanks for the paper!

Jeremy, why does the colour become so much more intense in the solid state than in solution, is this a feature of n to pi* transitions?  One compound is yellow with a faint green colour that becomes increasingly deeper green on concentration and crystallises out as a dark green solid (bit like crystalline fairly liquid!).  I have other equally deep green compounds (tetraazapentacenes) which on dilution remain the same colour, I think that with the azapentacenes the colour originated from pi to pi* transition. 

Well it might be worth doing concentration-dependent UV/vis to see if the colour arises from monomers or whether you have some degree of aggregation in solution. If you plot absorbance vs concentration you should get a linear dependence for monomer but a squared dependence if dimers are being formed.

Jeremy, but then I should expect to see the "aggregation" in the solid state. I got sheets separated by about 3.4 A and a whole load of Cl...Cl and Cl...S contacts at around 3.5 A, pretty much the sum of the vdw radii. Some of these contacts must not be so innocent. I guess its off to the UV/vis then. Tks.

Hi Tarunpreet, I think n-pi* do occur in carbonyls but the energy of that transition is higher so does not fall in the visible. As usual with spectroscopy there is the allowedness of the transition and its energy. The colour comes from the energy gap. If this is a discrete molecular phenomenon then it should not change with concentration (exept to increase in intensity due ot Beer-Lambert law) but if it is due to aggregation then new bands should be observed.

I was reading the conversation and would like to ask Jeremy the reason (the mathematical expression?) for the squared dependence of absorbance vs concentration when dimers are forrmed. Thank you

Hi Tarunpleet, sulfur has d orbitals that oxygen doesn't have. It is easy for sulfur to host electrons thru resonance and those electrons can circulate freely within those d energy levels which are not all exactly the same, thus producing colors. I believe it is the same phenomenon for many of the metal complexes which produce beautiful colors. For more info, try the Alexei, a 60' inorg chem russian book that was translated in various languages.

Dear Stefania,

If there is an equilibrium between monomer and dimer due to aggregation then we can write:

2 M   D

for which the equilibrium constant K can be written K = [D]/[[M]^2

So [D] = K[M]^2

Assuming both monomer and dimer obey Beer-Lambert law A = ecl, then the absorbance of the dimer band should follow [monomer]^2

Hope that helps.  

Amazing conversation going. Tarunpreet! I have handled compounds where I have seen solvent modifying the energy gap and solution of a colorless solid appearing colored. I have always associated color with extended conjugation and lowering of energy gap. Jeremy taught something great about poorer pi bonding between C=S due to bigger size I guess?  

Yes! Thank you

The strength of covalent bonding depends upon energy match, (symmetry match) and spatial overlap.  Pi bonding between elements of the second series (C,N,O) is strong because of good energy match and overlap between 2p orbitals. Conversely the 3p orbitals are more radially diffuse so although the energy match is good the amount of overlap of electron density is poorer and the tendency to form pi bonds is reduced. This is why we see lots of simple gaseous molecules such as N2, CO, CO2, O3 and so forth in which there is pi-bonding. When pi bonding is inefficient the atoms prefer sigma bonding (better overlap) thus nitrogen form N2 whereas phosphorus forms P4 (with just sigma bonds holding it together). Similarly O2 is a dimer with pi bonding whereas sulfur forms S8 with just a sigma framework. When we go to C=S we have bigger problems because the C 2p and S 3p orbitals have a poor energy match and spatial overlap, so the pi bonding is really weak. In many cases such bonding is better represented in the zwitterionic form, for example in DMSO we can write Me2S(+)-O(-) or in triphenylphosphine oxide, Ph3P(+)-O(-), indicating little pi character.

Amazing conversation going. Tarunpreet! I have handled compounds where I have seen solvent modifying the energy gap and solution of a colorless solid appearing colored. I have always associated color with extended conjugation and lowering of energy gap. Jeremy taught something great about poorer pi bonding between C=S due to bigger size I guess?  

2.How can we predict the color of a complex of extensively conjugated organic compound with a metal 

Can we predict how much change will be there in absorbing wavelength with every increase in conjugation