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.
