Theoretical principles


Many molecules absorb ultraviolet or visible light. The absorbance of a solution increases as attenuation of the beam increases. Absorbance is directly proportional to the path length, b, and the concentration, c, of the absorbing species. Beer's Law states that

A = ebc, where e is a constant of proportionality, called the absorbtivity.

Different molecules absorb radiation of different wavelengths. An absorption spectrum will show a number of absorption bands corresponding to structural groups within the molecule. For example, the absorption that is observed in the UV region for the carbonyl group in acetone is of the same wavelength as the absorption from the carbonyl group in diethyl ketone.

For a comprehensive discussion of Beer's Law, click here

Electronic transitions

The absorption of UV or visible radiation corresponds to the excitation of outer electrons. There are three types of electronic transition which can be considered;
  1. Transitions involving p, s, and n electrons
  2. Transitions involving charge-transfer electrons
  3. Transitions involving d and f electrons (not covered in this Unit)

When an atom or molecule absorbs energy, electrons are promoted from their ground state to an excited state. In a molecule, the atoms can rotate and vibrate with respect to each other. These vibrations and rotations also have discrete energy levels, which can be considered as being packed on top of each electronic level.

Absorbing species containing p, s, and n electrons

Absorption of ultraviolet and visible radiation in organic molecules is restricted to certain functional groups (chromophores) that contain valence electrons of low excitation energy. The spectrum of a molecule containing these chromophores is complex. This is because the superposition of rotational and vibrational transitions on the electronic transitions gives a combination of overlapping lines. This appears as a continuous absorption band.

Possible electronic transitions of p, s, and n electrons are;

s ® s* Transitions

An electron in a bonding s orbital is excited to the corresponding antibonding orbital. The energy required is large. For example, methane (which has only C-H bonds, and can only undergo s ® s* transitions) shows an absorbance maximum at 125 nm. Absorption maxima due to s ® s* transitions are not seen in typical UV-Vis. spectra (200 - 700 nm)

n ® s* Transitions

Saturated compounds containing atoms with lone pairs (non-bonding electrons) are capable of n ® s* transitions. These transitions usually need less energy than s ® s * transitions. They can be initiated by light whose wavelength is in the range 150 - 250 nm. The number of organic functional groups with n ® s* peaks in the UV region is small.

n ® p* and p ® p* Transitions

Most absorption spectroscopy of organic compounds is based on transitions of n or p electrons to the p* excited state. This is because the absorption peaks for these transitions fall in an experimentally convenient region of the spectrum (200 - 700 nm). These transitions need an unsaturated group in the molecule to provide the p electrons.

Molar absorbtivities from n ® p* transitions are relatively low, and range from 10 to100 L mol-1 cm-1 . p ® p* transitions normally give molar absorbtivities between 1000 and 10,000 L mol-1 cm-1 .

The solvent in which the absorbing species is dissolved also has an effect on the spectrum of the species. Peaks resulting from n ® p* transitions are shifted to shorter wavelengths (blue shift) with increasing solvent polarity. This arises from increased solvation of the lone pair, which lowers the energy of the n orbital. Often (but not always), the reverse (i.e. red shift) is seen for p ® p* transitions. This is caused by attractive polarisation forces between the solvent and the absorber, which lower the energy levels of both the excited and unexcited states. This effect is greater for the excited state, and so the energy difference between the excited and unexcited states is slightly reduced - resulting in a small red shift. This effect also influences n ® p* transitions but is overshadowed by the blue shift resulting from solvation of lone pairs.

Charge - Transfer Absorption

Many inorganic species show charge-transfer absorption and are called charge-transfer complexes. For a complex to demonstrate charge-transfer behaviour, one of its components must have electron donating properties and another component must be able to accept electrons. Absorption of radiation then involves the transfer of an electron from the donor to an orbital associated with the acceptor.

Molar absorbtivities from charge-transfer absorption are large (greater that 10,000 L mol-1 cm-1).

Review your learning

You should now be aware of why molecules absorb radiation in the UV and visible light regions, and why absorption spectra look the way they do.

After a short quiz, we will be considering the practical aspects of UV - Vis. spectroscopy, and looking at the instrumentation needed to perform this technique.

Beers Law - Quiz
UV-Vis. Absorption Spectroscopy - Theoretical Principles
UV-Vis. Absorption Spectroscopy - Theoretical Principles Quiz

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