What are Sigma and Pi bonds?
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:
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:
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.
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.
Sigma (σ) Bonding:
To understand Sigma bonding let us look at the simple molecule of methane (CH4).
We may all be familiar with drawing methane using electron dot diagrams, which would look something like this:
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.
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.
Fig 3: Energy diagram for Carbon
So why is methane written as CH4 and not CH2? Well the answer to this lies in something know as hybridization.
Forming the bonds
Now that we have hybridized the s and p orbitals of carbon to form four identical sp3 hybrid orbitals, it is time to bring in the Hydrogens that we ignored earlier. It is easy to see that the the four Hydrogens that will bond with the carbon all have a single 1s orbital with a single unpaired electron in each. This makes it very easy for it to bond with the carbon.At this point it is important that we notice how the 1s orbital of each of the four hydrogen comes together with the sp3 orbital of the central carbon. The two types of orbitals overlap in an end-to-end manner and form four single bonds which are referred to as sigma bonds giving us our methane molecule. Now remember the energy that the carbon atom gained to promote one of its electrons from the 2s to the 2pz orbital during hybridisation? Well once the carbon bonds with the hydrogens to form the CH4 molecule, it loses far more energy compared to this gain which eventually makes the molecule very stable and this is what is would look like:
Now lets try to recap what we have learnt using the ethane C2H6 molecule.
Fig 6: Sigma & Pi bonding in Ethane molecule. Source: chem1 virtual textbook home page
pi (π) bonds
To understand pi bonding lets have a look at the simple molecule of Ethene C2H4.
You may have drawn the ethane molecule many times in your classrooms and we are all aware of how the atoms and bonds are drawn to represent this molecule. Usually is would look something like this:
Fig 7: Structure of Ethene
In the case of Ethene, there is a difference from methane or ethane, because each carbon is only joining to three other atoms rather than four. When the carbon atoms hybridise their outer orbitals before forming bonds, this time they only hybridise three of the orbitals rather than all four. They use the 2s electron and two of the 2p electrons, but leave the other 2p electron unchanged.
The two carbon atoms and four hydrogen atoms would look like this before they joined together:
Fig 9: 1s and sp2 hybrid orbitals
The various atomic orbitals which are pointing towards each other now merge to give molecular orbitals, each containing a bonding pair of electrons. These are sigma bonds - just like those formed by end-to-end overlap of atomic orbitals that we saw in methane and ethane
The p orbitals on each carbon aren't pointing towards each other, are overlapping sideways.
This sideways overlap also creates a molecular orbital, but of a different kind. In this one the electrons aren't held on the line between the two nuclei, but above and below the plane of the molecule. A bond formed in this way is called a pi bond. Look at this illustration and notice how the orbitals have arranged themselves to form the pi bonds:
Clark, J 2010, Helping you to understand Chemistry. Chemguide, Retrieved September 12, 2010 from, http://www.chemguide.co.uk/index.html#top