Chemistry, Third edition

ENERGY LEVELS OF ATOMS AND MOLECULES

Energy levels of atoms and molecules

We now look at the types of energy possessed by molecules in more detail (Fig. 20.4).

Atoms and molecules possess electronic energy– energy due to the electrons in the

atomic or molecular orbitals. In addition to electronic energy, molecules (but not

isolated atoms) also possess vibrational energy– energy which causes the atoms in

molecules to vibrate.

Most of the energy of an atom or molecule is in the form of electronic energy. Mak-

ing an analogy with money, electronic energy is ‘big money’ (pounds), whereas vibra-

tional energy is ‘small money’ (pence).

The quantum theory only allows the electronic and vibrational energy of

molecules to hold certain values. The energy of a molecule may be looked upon as a

‘ladder’ of allowed electronic energy levels or ‘states’ (labelled E 0 ,E 1... ) with sub-lev-

els according to the vibrational energy (labelled v 0 , v 1 , v 2... ) possessed by the

molecule. For example, a molecule in its second electronic energy level and its third

vibrational level would be labelled E 1 ,v 2.

The lowest energy level of a molecule or atom is called its ground state, whereas

higher levels are referred to as excited states. The making of excited states (by photon

absorption or by raising the temperature of the sample) is called excitation. At room

temperature nearly all atoms and molecules lie in their ground electronic and (for

molecules) in their ground vibrational energy levels, i.e. they lie at E 0 and v 0. This

means that, except at very high temperatures, molecular or atomic transitions involv-

ing absorption always start from the lowest energy level.

Transitions involving changes in electronic energy result in electrons moving

from one orbital to another, the electrons being promoted to a higher energy orbital

in absorption and demoted to a lower energy orbital in emission. The patterns

observed in a spectrometerand which are caused by these transitions are referred to as

electronic spectra. The transitions involve so much energy (typically 100 kJ per

mole of absorbing atoms or molecules) that photons of ultraviolet or visible light are

involved, and so the spectra are also called UV–visible spectra. In Fig. 20.4, transition

‘a’ involves the absorption of UV–visible light– this is what happens when a molecule

of dye absorbs visible light. Transition ‘c’ involves the emission of UV–visible light.

Transitions in which the vibrational energy (only) of molecules is changed

involve photons of infrared light, and the resulting spectroscopic patterns are

referred to as infrared spectra. The energy jumps (E) involved in infrared spectra

are smaller than those involved in UV–visible spectra, and are typically 10 kJ per

20.2

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Fig. 20.4Electronic and vibrational

energy levels of an isolated molecule. This

diagram is a more detailed version of Fig.

20.3. The transitions labelled a and c

involve the absorption and emission

(respectively) of ultraviolet or visible light.

The transitions labelled b and d involve

the absorption and emission (respectively)

of infrared light.