Does delocalisation always lead to molecular stabilisation?


Does delocalisation always lead to molecular stabilisation?

30 JULY 2021

Most bonds we are familiar with still have a localised nature, with the highest electron density occurring within the internuclei region, and negligible electron density far away from the internuclear region. And yet, such is the beauty of chemistry, that there always are exceptions, and delocalised bonding patterns in molecules like benzene is an example known to all high school students. The nature of such bonding is that no longer is the bulk of the electron density located within the internuclear region of two atoms, but the electron density instead reveals an increased dispersion of the electron density over three or more atoms. And this delocalisation at the molecular level leads to observable properties “surprising” to those only acquainted with a localised picture of bonding — the electrical conductivity of graphite is a good example, as well as the brilliant range of colours of small organic compounds with significant conjugation. We are often taught that delocalisation of electrons stabilises a molecule — but is this always true?

Delocalisation can confer stability to charged chemical species through the dispersal of the charge across multiple atoms, increasing the stability of the compound. An example is the allylic cation, shown below.

In simple terms, the pi electron cloud of the C=C bond overlaps with the empty p orbital “holding” the positive charge, allowing for the delocalisation of the positive charge over the entire molecule, stabilising it. Perhaps another way to think of it would be a donation of electron density from the pi bond to the empty p orbital on the positively charged carbon, stabilising the molecule through electrostatic attraction. Whether one thinks of it as the positive charge being dispersed or electrons being donated, the result is the same. The positive charge is now shared by the two terminal carbon atoms, and the molecule is now better represented in the form of two resonance structures:

Likewise for the allyl carbanion, the delocalisation of the negative charge across the molecule helps to reduce electrostatic repulsion between the electrons, stabilising the carbanion.

It is easy to see how delocalisation can stabilise charged molecules. It is less straightforward to see how delocalisation does in fact stabilise radical species. In school, we are taught the simplified version that radical species are “electron-deficient” and would benefit from the donation of electron density through delocalisation. However, that is not a very convincing argument, as the molecule has an overall neutral charge. Perhaps a better understanding can be reached if one considers the delocalisation through the lens of molecular orbital theory. In this case, we can consider the overlap of the three atomic p orbitals to form 3 molecular orbitals.

As shown in the diagram above, one can immediately see the energetic stabilisation conferred by the interaction of all three p atomic orbitals, compared to the case where only two of them form a pi bond while the third holds the unpaired electron.

Molecular orbital theory further allows us to understand how delocalisation can confer special stability to cyclic compounds, termed as aromaticity. Considering only the six p atomic orbitals of benzene, and combining these orbitals to form six molecular orbitals, one gets an MO diagram consisting of 3 bonding orbitals and 3 antibonding orbitals.

One finds that by having 6 electrons in the pi system, one can fill up the bonding orbitals completely. This leads to an extremely stable configuration, explaining why benzene does not undergo addition reactions, as the conversion of any carbon from sp² to sp³ destroys the symmetrical pi system and raises the energy of the system. One can further prove Huckel’s Rule, that a planar ring system would experience this stabilising aromatic effect when the pi system contains (4n + 2) electrons. Hence for such ring systems, delocalisation does contribute greatly to energetic stability.

However, sometimes a molecular system without electron delocalisation can be more stable than a delocalised system when the delocalisation results in an antiaromatic ring. The reasons for anti-aromatic compounds being unstable is complex, but according to this source, the reason for this is the existence of partially filled delocalised orbitals being susceptible to chemical attack. Consider the cyclopropenyl anion. If its geometry allowed delocalisation to occur, it would be antiaromatic, with the corresponding molecular orbital diagram shown below. Not only are there two partially filled orbitals, they are also antibonding orbitals, hence being more energetically unstable than nonbonding orbitals.

As such, the p orbital holding the lone pair of electrons in the cyclopropenyl anion exists out of the plane of the molecule and there is no delocalisation of the pi electrons, as this results in greater stability of the molecule.

In addition, the fulfilment of planarity to achieve delocalisation in a ring structure could result in decreased stability of the molecule. A good example would be [10]annulene. If [10]annulene were to exist as a planar molecule, the C-C-C bond angles would be 144 degrees, significantly different from the ideal bond angle of 120 degrees for sp2 orbitals, resulting in less effective orbital overlap. Hence, [10]annulene sacrifices the stability offered by aromaticity, having a non-planar most stable conformer which minimises ring strain.

Evaluating, it is clear that the energy level of a molecule is affected by many electronic factors, such as steric repulsion, effectiveness of orbital overlap, as well as many other factors beyond our usual consideration. Hence, it is clear that while delocalisation could contribute to the stability of a species, it would be premature to assume that delocalisation of electrons in all cases would lead to a stabilisation of a chemical species, given the examples of antiaromaticity and nonaromatic annulenes described above.

Ian Tay

Author | Editor