The electron movements behind UV/Visible spectroscopy
by James Everitt.
Electrons reside in orbitals around an atom. The most basic of these are the s and p orbital. Every electron shell in an atom contains one spherical shaped s-orbital which when it combines with another to form a bond forms what is known as a σ-bond which is an oval shaped shared region for the 2 bonding electrons. A p-orbital is a lobe shape which exists in each of the 3 plains (x,y and z). Each shell from the second contains 3 p-orbitals which can form 1 σ-bond and 2 π-bonds. These π-bonds are characteristic of double bonds and are the electron regions that interact with electromagnetic radiation and absorb some of the light.
In UV/Visible spectroscopy both non-bonding and π-bonding orbitals are affected. The electrons within these regions are excited to an antibonding state called π*. A non-bonding orbital has a higher energy due to its greater density which means it is often the orbital that is excited by visible light (because visible light has a lower energy than UV light). The π-bonding orbitals are excited by the UV light giving the UV absorption.
π-bonding orbitals are often more common and occupy much more space that non-bonding orbitals meaning the π-bonding orbitals absorb much more that the non-bonding giving them a much larger peak on the absorption graph.
Conjugation is where a molecule contains multiple double bonds within its structure. This affects the UV/Visible spectroscopy by brining down the π* energy level meaning longer wavelengths are absorbed. This means for large molecules with many double bonds this absorption happens almost entirely in the visible spectrum. This leads to the colouring of large molecules such as β-Carotene which gives carrots and tomato’s their colour and Melanin which is the pigment that colours your hair and skin
Melanin (Eumelanin) β-Carotene