Biochemical Techniques and Instrumentation

1.1.2. Interaction of EMR with matter

Light is electromagnetic radiation (EMR) that exhibit discrete packets of energy called photons. Electromagnetic spectrum consists of wavelength from the order of meters ie. radio waves to less than 1 A^o ie. gamma rays. The wavelength of visible light is in a narrow range of between 4000 A^o to 8000 A^o (400 - 800nm). The transitions that occur due to the interaction of electromagnetic radiation (photons) with matter is a quantum phenomenon.  The interaction is dependent upon the properties of the radiation and the appropriate structural components of the molecules. The photons interact with matter by transferring its energy.

Atoms and molecules occupy distinct energy levels. They have unique energy configurations in terms of electron configurations, vibrational/rotational levels, etc. The Jablonski diagram gives the electronic and vibrational states of the molecule. Electrons in the atoms or molecules are distributed between several energy levels. Electrons principally reside in the lowest level or ground state (So). When energy is absorbed, it is promoted to a higher level or excited state (S1). The energy from electromagnetic radiation gives rise to this absorption spectrum. The molecule will also be in an excited vibrational and rotational state. The molecule will subsequent relax into the vibrational ground state of the first electronic excited state. The electron can then revert back to the electronic ground state. This will be accompanied by the emission of heat in case of non-fluorescent molecules.

The plot of absorption against wavelength is called absorption spectrum. Single atoms give line spectra, while molecules give band spectra due to different kinds of energy levels. Molecular spectra are molecule specific due to the unique vibrational states.

The energy diagram of molecules is more complex as they possess more than two atoms. The different atomic orbitals combine to yield molecular orbitals. When two atomic orbitals combine they form both a low energy bonding molecular orbital, and a high energy antibonding molecular orbital (*). The s and p atomic orbitals overlap in different ways. The head-on overlap of two s or two p atomic orbitals forms σ bond and the parallel overlap of two p atomic orbitals forms π bond. Single bonds are usually σ bonds, where as double bonds contain one σ bond and one π bond. Each of these molecular orbitals represents a different energy level.

Each atom in an organic molecule in its ground state or low energy state will contain bonding electrons in σ and π molecular orbitals, and outer, nonbonding unshared electrons 'n'. Molecular orbitals generally fall into one of the five different classes: s orbitals combine to the σ and the anti-bonding σ* orbitals. Some p orbitals combine to the bonding π and the anti-bonding π* orbitials. Other p orbitals combine to form non-bonding n orbitals. The population of bonding orbitals strengthens a chemical bond and the population of anti-bonding orbitals weakens a chemical bond.