BNG_Sigma and Pi Bonds Lesson

Sigma and Pi Bonds

There are a couple more terms that are used to describe covalent bonds. These terms, sigma and pi bonds, refer to the area around the nucleus where the bond is formed.

When two orbitals overlap directly between two bonded atoms they form a molecular orbital called a sigma bond ( σ bond). Sometimes this area of space is called the inter-nuclear axis. Any bond of this kind, whether formed from the overlap of s orbitals, p orbitals, or hybrid orbitals, is called a sigma bond. Sigma bonds are the strongest type of covalent bonds due to the direct overlap of orbitals

There is another way that p orbitals can overlap. Instead of overlapping end to end, as shown in (b) above, they could overlap side by side. This type of bond is called a pi-bond (π bond). Pi-bonds are defined as the parallel overlap of p-orbitals. This produces a bond in which the electron density is divided between two separate regions that lie on opposite sides of an imaginary line joining the two nuclei.    

Look at this picture of a π bond. You might think that you are looking at two bonds, but you are not. A π bond has two parts. It takes both of them to equal one π bond.

Remember that sigma and pi bonds refer to the relationship in space around the nucleus.   They don't indicate the type of orbital.   The overlap is stronger in sigma than pi bonds, which is reflected in sigma bonds having larger bond energy than pi bonds.

Double and Triple Bonds

When we draw the Lewis structure for a double or a triple-bond, like ethene or ethyne below, it seems as if we are implying that either four or six electrons are directly between these two atoms.

We know that an orbital (atomic or hybrid) can only contain two electrons. So, how is a double or triple bond able to form? It is the formation of π bonds that allows atoms to form double and triple bonds. Remember that π bonds form above and below the axis and that σ bonds form between the nuclei.

Here is another drawing of ethene:

So, a double bond has one sigma bond and one pi bond. This is true of all double bonds.  

The next question we need to address is which orbitals formed each of these bonds.

• The pi bond is formed by the overlap of two p orbitals. All pi bonds are formed from p orbitals.
• Determining the orbitals involved in the sigma bond is trickier. Remember that sigma bonds can form from the overlap of any type of orbital, hybrid or not. As long as the overlap occurs in the area between the two nuclei it is a sigma bond. In order to determine the orbital type that is participating, we need to refer VSEPR. Look at each C independently. Each C has three bonding clouds and no lone pairs. This means that the shape is trigonal planar. Trigonal planar has bond angles of 120^o and is sp² hybridized. So, the sigma bond is formed from the overlap of two sp² orbitals. The image to the right shows the sp² orbitals in blue and the p orbitals in green.

Acetylene (ethyne) $H-C \equiv C-H$  forms in a similar manner. The triple bond consists of one sigma bond and two pi bonds. The first pi bond is in the same location as shown in ethene above. The second pi bond is in the third dimension, coming out of the page, as shown below.

To summarize:

• The first covalent bond between atoms is always a sigma bond.
• All single bonds are sigma bonds.
• To determine the orbitals involved in sigma bonds, look at the VSEPR geometry and matching hybridization.
• The second or third bond formed between two atoms is a pi bond
• Pi bonds are always formed from two unhybridized p orbitals.

Remember to work on the module practice problems as you complete each section of content.  

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