Our familiar model of covalent bonding is largely based on the Lewis model, with atoms sharing pairs of bonding electrons and everyone obeying the octet rule. It is often very successful in explaining the structures of molecules, and the bonding that gives rise to those structures, but there are many exceptions.
Try drawing out a dot and cross diagram for nitrogen monoxide. Either you can go for a double bond, in which case nitrogen has only 7 electrons, or you can try a triple bond, but then oxygen has 9 electrons. In phosphorus pentafluoride, the phosphorus has 5 bonding pairs of electrons surrounding it (we explained this earlier by invoking the idea of an expanded octet). Or take oxygen. The Lewis structure says that there is a double bond between the two oxygen atoms and each atom also has two lone pairs of electrons. The trouble is that liquid oxygen is paramagnetic, which can only be if there are unpaired electrons in the structure.
There are two modern theories of bonding – valence bond theory and molecular orbital theory. Each has its strengths and weaknesses and, like the Lewis model, is useful in different situations.
Both are based on the idea that atomic orbitals interact with each other when a covalent bond is formed between the two atoms, but in order to overcome the problem that we can’t solve Schrödinger equation for atoms other than hydrogen, the two theories make differing approximations and assumptions.
This is extension work – unless you are following the Cambridge Pre-U specification – but an overview of the two theories and will give you a more complete understanding of how the concept of bonding has developed in modern times. If you are thinking about studying Chemistry, Natural Science or a related degree at university, then I strongly recommend you take a look.
Valence bond theory
Valence bond theory is particularly useful for describing bonding in organic carbon-based molecules.
E.g. H2
- the 1s orbital on each hydrogen atom interact to form a 𝜎-bonding orbital, positioned between the two hydrogen atoms and the 1s electrons are placed in this orbital
- in valence bond theory the bonds have both ionic (H+ H–) and covalent character – how each is weighted depends on the molecule. The real structure is a mixture of the different forms
E.g. CH4
- in valence bond theory the atoms in molecules can change the shape and energy of their atomic orbitals in a way that increases the opportunities for bonding – this is known as hybridisation
- carbon hybridises the 2s and 2p orbitals to form four sp3 hybrid orbitals which have the same energy as the original s and p orbitals had between them
- in methane each sp3 hybrid orbital interacts with a 1s orbital from a hydrogen atom to form a 𝜎-bonding and 𝜎*-antibonding orbital – the bonding electrons are placed in these new orbitals
E.g. C2H4
- in ethene the 2s, 2px and 2py orbitals on each carbon atom hybridise to give three sp2 hybrid orbitals
- one of these sp2 hybrid orbitals interacts with an sp2 hybrid orbital from another carbon atom, and the other two sp2 hybrid orbitals interact with 1s orbitals from hydrogen atoms – we now have three 𝜎-bonding orbitals each holding a pair of bonding electrons
- the unhybridised 2pz atomic orbitals on each carbon atom interact to give a ∏-bonding orbital, containing a pair of electrons, with electron density extending above and below the plane of the 𝜎 bonds
