Professor Dave here, let’s get a closer look at covalent bonds.
We know a lot about chemical bonds, both ionic and covalent.
But let’s take a closer look at the covalent bond now.
Whereas an ionic bond is an electrostatic interaction between formally charged ions,
a covalent bond involves two atoms, typically two non-metals, sharing two electrons, such
that each electron is interacting with both of the nuclei.
Now it’s time to get a bit more quantitative than this, with the help of valence bond theory.
Let’s look at two hydrogen atoms.
Given that these have precisely the same ionization energy and electron affinity, electron transfer
will not occur, so an ionic bond is not a possibility.
Now let’s take a look at this diagram, which depicts the potential energy of the two hydrogen
atoms as a function of the distance between them.
Beginning at the right edge, the atoms are far apart and not interacting, so the potential
energy is zero.
Let’s see what happens as they move closer together.
As they approach one another, their 1s orbitals begin to overlap, allowing each electron to
interact with the proton of the other hydrogen atom.
As this continues, the potential energy of the system gets lower and lower, until an
ideal internuclear distance is found, the configuration with the lowest possible energy.
This is the ideal distance between the two protons and therefore the precise bond length
for a molecule of hydrogen.
This decrease in potential energy occurs because each electron is interacting with both protons,
and this is a stabilizing and therefore energetically favorable situation.
However, if we continue to push the protons together, things become less favorable, due
to cumulative proton-proton repulsion and electron-electron repulsion, which will quickly
begin to outweigh any proton-electron attraction.
This is sort of like if we were to compress a spring beyond its equilibrium position,
and this is why we see the curve rise steeply to the left.
So as we said, it is this phenomenon that reveals the internuclear distance that offers
the lowest potential energy and therefore the ideal covalent bond length for the hydrogen
atoms, which for this case will happen to be 74 picometers, where a picometer is a trillionth
of a meter.
This also shows why covalent bond formation is always an exothermic process, because there
will always be a decrease in potential energy as a result of this kind of orbital overlap,
and that difference in energy must be released to the environment.
This is also the energy that is required to break the bond, and it will be equal to 7.24
x 10^-19 joules for one hydrogen molecule.
That’s a very tiny energy, but multiplying by Avogadro’s number gives us the energy
required to break the covalent bonds in a mole of hydrogen molecules, and that will
be 4.36 x 10^5 joules, which becomes quite significant.
In this way, we can look at any covalent bond in a slightly more sophisticated way.
Sigma bonds involve direct orbital overlap, and this can be between atomic orbitals, or
between two hybrid orbitals, be they sp, sp2, sp3, sp3d, or sp3d2, the ones we learned about
from VSEPR theory.
They can even be between an atomic orbital and a hybrid orbital.
Pi bonds, on the other hand, involve lateral orbital overlap between unhybridized p orbitals.
But no matter what the case, for any covalent bond, there is some kind of orbital overlap,
and the pair of electrons will inhabit the overlapping region.
It is the mutual attraction between both electrons and both atomic nuclei that facilitates this link.
And the more significantly these orbitals overlap, the stronger the covalent bond will
be, and the lower the potential energy will be for the system.
This is why a shorter bond is a stronger bond, because it signifies greater orbital overlap.
The same goes for direct overlap versus angled overlap, direct overlap is more significant,
and thus results in a stronger bond.
We can assign bond energies to all types of covalent bonds, and they will all be different,
because they all involve atoms of different elements, with different electronegativities.
Different orbitals will overlap in different ways, so different amounts of energy will
be released upon their formation.
Hopefully, we now have a slightly better understanding of covalent bonds.