The energy of molecules is quantised just as it is in atoms.
When a molecule absorbs a photon of electromagnetic radiation, it jumps from one molecular energy level to another as long as the energy of the photon corresponds exactly to the difference between the two energy levels.
Consider a molecule of HBr. If we shine electromagnetic radiation at a gaseous sample with energy in the range of 1 x 10-20 J to 1 x 10-19 J the HBr molecule will absorb a photon with a specific energy or frequency and it will move from a lower to a higher vibrational energy level. The molecule will vibrate more strongly.
The photon with the specific energy or frequency of radiation needed to promote a molecule of CO2 from a lower to a higher vibrational energy level will be different to that of HBr, because ΔE (the difference between the vibrational energy levels) is different. ΔE is specific to the bond and the molecule.
Molecular spectroscopy covers a range of techniques in which we investigate how different energies or frequencies of electromagnetic radiation interact with molecules. This gives us useful information about how the atoms in molecules are arranged, and the bond strengths, lengths and bond angles in a molecule.
High energy ultraviolet radiation
High energy uv radiation has an energy similar the bond energies in molecules. When molecules absorb such high frequency radiation covalent bonds are broken and radicals are formed.
This photodissociation reaction is the initiation step in photochemical reactions such as those involving ozone in the stratosphere.
Low energy ultraviolet and visible radiation
When molecules absorb low energy uv or visible radiation electrons are promotes from a lower energy molecular orbital to a higher energy molecular orbital that is either unoccupied or partially occupied.
This is an electronic transition and explains why some molecules are coloured and some are colourless. If a molecule is colourless ΔE corresponds to the absorption of photons of low energy uv light which is not detected by our eyes. In coloured molecules ΔE corresponds to the absorption of photons of visible radiation and we ‘see’ the molecule’s colour in terms of all the visible light that is transmitted or reflected.
- in transition metal complexes absorption of visible light leads to the transition of electrons either between d-orbitals (this is why Cu2+(aq) is a pale blue solution) or between an orbital in the ligand and an orbital in a metal ion (this is why the manganate(VII) ion, MnO4–, is deep purple).
- in organic molecules absorption of low energy uv or visible radiation causes the transition of electrons between ∏- orbitals in conjugated systems (a conjugated system is where a molecule contains alternating single C-C and double C=C, N=N or C=O bonds).
E.g. plant leaves mostly appear green because they contain chlorophyll pigments, including chlorophyll a and chlorophyll b.
Chlorophyll a absorbs violet and orange light and chlorophyll b absorbs mostly blue and yellow light. Neither pigment absorbs green light which is reflected, hence leaves appear green coloured.
Note that the very small structural change in the conjugated system between the two pigments (simply swapping a methyl group for an aldehyde functional group, CHO) has a significant effect on the uv-vis absorption spectrum, as indicated by the red arrows in the diagram above.
The conjugated system is also called the chromophore – essentially this is the part of the molecule that is responsible for the absorption of visible light that results in the molecule being coloured.
Infrared radiation
The absorption of infrared radiation causes covalent bonds to vibrate more strongly – the bond lengths and bond angles of the molecule change.
This is the basis of infrared spectroscopy which is primarily used to identify the functional groups in molecules, but it also helps us explain why we feel warm if we are standing in the sun – the molecules in our skin absorb specific frequencies of IR radiation which causes them to vibrate more strongly and giving them more kinetic energy which we sense as warmth.
Microwave radiation
The energy or frequency of microwave radiation corresponds to transitions between rotational energy levels in molecules.
Rotational spectroscopy can be used to determine bond strengths but we can also use it to monitor the concentrations of ClO• radicals in the stratosphere, helping us to understand the role of CFCs in ozone depletion.
Questions
- Calculate the energy of a photon of infrared radiation with a frequency of 5.20 x 1013Hz, given that Planck’s constant is 6.63 x 10-34 JHz-1.
- Explain why greenhouse gases only absorb certain frequencies of the IR radiation emitted by the Earth and suggest why this leads to a warming of the troposphere.
- Photochemical smog is the result of a series reactions involving primary and secondary pollutants. In one of these reactions nitrogen(IV) oxide is photolysed to form nitrogen(II) oxide and an oxygen atom.
NO2• + ℎ𝜈 ⇾ NO• + O ΔrH⦵ = +214 kJ mol-1
Calculate the frequency of visible radiation required to break a single N-O bond in a molecule of NO2.
Answers
- 3.45 x 10-20 J
2. Greenhouse gas molecules absorb specific frequencies of IR radiation which correspond to the transition / ΔE between a lower vibrational energy level and a higher one. The molecules vibrate more strongly / their vibrational energy increases and they have more kinetic energy which is transferred to other molecules in the air through the increased number of collisions between them.
3. Energy required to break one N-O bond in J = 214000 / (6.02 x 1023) = 3.55 x 10-19 J
𝜈 = E / h = (3.55 x 10-19) / (6.63 x 10-34) = 5.36 x 1014 Hz