16) Discuss the shapes of s. p & d orbitals. Explain the formation of sigma and pie bond with orbital
overlapping!
Answers
Answer:
Many of us are already aware of the definition of a sigma bond from our teachers, text books or from many of the websites online. However, if you are still not aware of what these two bonds are, then here is a basic definition of the two:
Sigma bond: A covalent bond resulting from the formation of a molecular orbital by the end-to-end overlap of atomic orbitals, denoted by the symbol σ.
Now have a look at this illustration to see how this end-to-end overlapping occures:

Fig 1: Formation of a Sigma bond
Misconception: many students in the Pacific may have this worng notion that a sigma
Pi bond: A covalent bond resulting from the formation of a molecular orbital by side-to-side overlap of atomic orbitals along a plane perpendicular to a line connecting the nuclei of the atoms, denoted by the symbol π.
Here's another illustration showing how the side-to-side overlapping occurs:

Fig 2: Formation of a Pi bond
It is important to note that different sources use different terms to define what a sigma and pi bond is. However, once examined carefully, it will be evident that they all try to explain the same thing.

Misconception: many students in the Pacific may have this wrong notion that a sigma bond is the result of the overlapping of s orbitals and a pi bond is the result of the overlapping of p orbitals because they may relate the 's' to 'sigma' and the 'p' to 'pi'. However, it is seen that sigma bonds can be formed by the overlapping of both the s and p orbitals and not just s orbital.
You may have noticed that in order to understand these definitions it is obvious that we must know what an s and p orbital is.
Please click here to learn more about Atomic Orbitals if you are unfamiliar with the concept.

Note: A single bond such as (C-H) has one sigma bond whereas a double (C=C) and triple (C≡C) bond has one sigma bond with remaining being pi bonds.
Bond type
No. of σ bond
No. of π bonds
Single (C-H)
1
0
Double (C=C)
1
1
Triple (C≡C)
1
2
Sigma (σ) Bonding:
To understand Sigma bonding let us look at the simple molecule of methane (CH4).
Methane, CH4
We may all be familiar with drawing methane using electron dot diagrams, which would look something like this:

Fig 3: Covalent bonding in Methane

Misconception: many students after drawing such electron dot diagrams fail to appreciate that in reality molecules exist as a 3D system and not as a two dimensional system as shown above. These diagrams are drawn for simplicity and should not be viewed as an exact representation of what a molecule looks like.
For now, let us ignore the Hydrogen and concentrate on the central Carbon atom. We know that it is the valence electrons that are responsible for covalent bonding and we must know the electron configuration of an element from the periodic table to know how many valence electron it has.
Please click here to learn more about Electron Configuration if you are unfamiliar with the concept.
Now, when we look at the carbon atom from our Methane, we see that its electron configuration is 1s2 2s2 2p2. However, from this electron configuration we can see that carbon has only two unpaired electrons (2p2) in its valence shell which can be used to form bonds with two hydrogen atoms. You can see this more clearly in the electrons-in-boxes notation below.