Here’s what this post is covering:
- Orbital Overlapping
As you know, electrons exist in ORBITALS.
The behaviour of covalent bonds can be explained in terms of orbitals of different atoms OVERLAPPING.
When orbitals overlap, their component electrons can exist within their combined space.
Thus, the electrons are considered to be SHARED between the parent atoms.
This is what COVALENT BONDING is.
There are 2 types of orbital overlapping: σ-bonds & π-bonds
Q: Why are they called sigma & pi?
What are Sigma Bonds (σ-bonds)?
Bonds caused by END-TO-END overlapping of orbitals.
These orbitals can be s orbitals and/or p orbitals.
Since s-orbitals are spherical, any bonding involving them is considered end-to-end:
For p-orbitals, if the overlapping involves these highlighted areas, it is considered end-to-end:
Examples of σ-bonds:
|Between s & s orbitals|
|Between p & p orbitals|
|Between s & p orbitals|
Characteristics of σ-bonds:
|Bond axis (the centre line of the bond) is along the internuclear axis (line joining the nuclei of both atoms)|
|Charge distribution is localised along the internuclear axis|
|Electrons are tightly bound & require high energy to remove|
|Thus, bond is very strong|
What are Pi Bonds (π-bonds)?
Bonds caused by SIDE-TO-SIDE overlapping of orbitals.
These orbitals can be p orbitals or d orbitals.
Examples of π-bonds:
|Between p & p orbitals|
|Between d & d orbitals|
|Between p & d orbitals|
Characteristics of π-bonds:
|The overlap of the orbitals (shown in blue) is 90 degrees to the internuclear axis of bonded atoms|
|Charge distribution is above & below the internuclear axis|
|Electrons are loosely bound & require low energy to remove|
|Thus, bond is weaker|
How can you know the type of bonding?
The FIRST overlap of 2 orbitals is ALWAYS end-to-end, since this is the simplest way for them to orient (with maximum concentration of charge).
However, if more than 1 pair of electrons needs to be shared, the next electrons must come from another orbital, which is PERPENDICULAR to the first set of orbitals (due to the way orbitals are arranged).
The first pair of electrons here come from px orbitals.
They naturally overlap end-to-end.
If it were a single bond, it would be a σ bond.
If I need to share another pair of electrons, they will be from either a py or pz orbital.
Both of which are 90 degrees from the px.
This double bond contains the original σ-bond, & a 2nd π-bond.
If I were to need another pair of shared electrons, they will be from the pz orbital.
Which is 90 degrees from both the px & py.
This triple bond contains a σ-bond, & 2 π-bonds.
A General Rule of Thumb is:
- Single-bond: always σ-bond
- Double-bond: always 1 σ-bond + 1 π-bond
- Triple-bond: always 1 σ-bond + 2 π-bond
|END-TO-END overlapping of orbitals||SIDE-TO-SIDE overlapping of orbitals|
|1 overlap in every bond (single, double, etc) is always a σ-bond…||…the remaining overlaps are π-bonds|