PRINCIPLE OF UV- VISIBLE SPECTROSCOPY

The Ultraviolet (uv) region extends from about 10 to 380 nm, but the most analytical useful region is from 200 – 380nm, called the near ultraviolet region. Below 200 nm, the air absorbs appreciably and so the instruments are operated under a vacuum; hence, this wavelength region is called the vacuum ultraviolet region. The visible region is the region of wavelengths that can be seen by the eye, that is, the light appears as a colour. It extends from the near ultraviolet region (380 nm) to about 780 nm. When an organic molecule absorbs UV- visible radiation, the energy from UV or visible light causes the outer electrons from a lower energy to be raised to a higher energy level, corresponding to an electronic transition. This transition of electrons is from molecular bonding orbital to the higher energy antibonding molecular orbital. According to the molecular orbital theory, the shared electron pair of a covalently bonded atoms may be thought of as occupying molecular orbitals (MO) which is of a lower energy and has a corresponding unoccupied orbitals called antibonding  molecular orbitals, these correspond to excited state energy levels (higher energy level). When the molecule is in the ground state, both electrons are paired in the lower- energy bonding orbital – this is the Highest Occupied Molecular Orbital (HOMO). The antibonding, in turn, is the Lowest Unoccupied Molecular Orbital (LUMO).

When organic molecules that are capable of absorbing UV visible radiation (chromophores) are exposed to the radiation at a wavelength with energy equal to the difference between the energy of the HOMO and LUMO (with energy equal to ΔE, the HOMO‐LUMO energy gap) this wavelength will be absorbed and the energy used to bump one of the electrons from the HOMO to the LUMO. The energy absorbed appear as absorption peaks at the wavelength it corresponds to on the UV – Visible spectrum. The extent of absorption of electromagnetic radiation corresponds to the concentration of the analyte through the application of Beer – Lambert law.

Molecules containing pi electrons or non-bonding electrons can absorb the energy in the form of ultraviolet or visible light to excite these electrons to higher anti- bonding molecular orbitals. The more easily excited the electrons (i.e. lower energy gap between the HOMO and the LUMO) the longer the wavelength of light it can absorb.