CHEM C3: Orbital Overlapping

Here’s what this post is covering:

  • Orbital Overlapping
    • Sigma-Bonds
    • Pi-Bonds

Let’s go!


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?
A: 


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:

Screenshot 2019-02-11 20.37.39

For p-orbitals, if the overlapping involves these highlighted areas, it is considered end-to-end:

Screenshot 2019-03-26 14.42.31.png

Examples of σ-bonds:

Between s & s orbitals http___www.ibchemistrytutors.com_wp-content_uploads_2015_06_sigma-bonding.png
Between p & p orbitals http___www.ibchemistrytutors.com_wp-content_uploads_2015_06_sigma-bonding-1.png
Between s & p orbitals http___www.ibchemistrytutors.com_wp-content_uploads_2015_06_sigma-bonding-2.png

Characteristics of σ-bonds:

Bond axis (the centre line of the bond) is along the internuclear axis (line joining the nuclei of both atoms) Sigma Bond | OChemPal
Charge distribution is localised along the internuclear axis Sigma bond - Wikipedia
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 File:Pi-Bond.svg - Wikimedia Commons
Between d & d orbitals Screenshot 2019-03-26 14.55.05.png
Between p & d orbitals Screenshot 2019-03-26 14.55.00.png

Characteristics of π-bonds:

The overlap of the orbitals (shown in blue) is 90 degrees to the internuclear axis of bonded atoms Screenshot 2019-03-26 15.04.06.png
Charge distribution is above & below the internuclear axis Screenshot 2019-03-26 15.05.45.png
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).

For example:
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.

http___www.ibchemistrytutors.com_wp-content_uploads_2015_06_sigma-bonding-1.png

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.

pi bond | Organic Chemistry

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.
8.1 Valence Bond Theory – Chemistry

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

In Summary,

σ-bonds π-bonds
END-TO-END overlapping of orbitals SIDE-TO-SIDE overlapping of orbitals
Chemical bonding and structure, IB DP Chemistry notes File:Pi-Bond.svg - Wikimedia Commons
  • Bond axis is along the internuclear axis
  • Charge distribution is localised along the internuclear axis
  • Electrons are tightly bound & require high energy to remove
  • Bond is very strong
  • The overlap of the orbitals is 90 degrees to the internuclear axis
  • Charge distribution is above & below the internuclear axis
  • Electrons are loosely bound & require low energy to remove
  • Bond is weaker
1 overlap in every bond (single, double, etc) is always a σ-bond… …the remaining overlaps are π-bonds

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