CN- Lewis Structure, Molecular Orbital Diagram, And, Bond Order

The Lewis structure of a cyanide [CN]– ion consists of a carbon (C) atom and a nitrogen (N) atom. The two atoms are connected via a triple covalent bond. There are a total of 3 bond pairs and 1 lone pair around both C and N atoms respectively in CN– lewis structure.

The 3 bond pairs are considered a single electron domain while determining the shape and/or geometry of the molecular ion.

In this article, we will discuss the Lewis structure of CN and CN+ as well, in addition to drawing the Lewis structure of CN–.

But let us first start with the simple steps given below to draw the Lewis dot structure of [CN]–.

Steps to draw the lewis dot structure for CN–

1. Count the total valence electrons in [CN]–

The Lewis dot structure of a molecule or a molecular ion is a simplified representation of all the valence electrons present in it. Therefore, the very first step while drawing the Lewis structure of the cyanide [CN]– ion is to count the total valence electrons present in its concerned elemental atoms.

There are 2 different elements present in [CN]– i.e., carbon (C) and nitrogen (N). When you search through the Periodic Table of elements, you will find that carbon is present in Group IV A so it has a total of 4 valence electrons in each atom. In contrast to that, nitrogen belongs to Group V A of the Periodic Table so it has a total of 5 valence electrons in each atom.

• Total number of valence electrons in carbon = 4
• Total number of valence electrons in nitrogen = 5

∴ [CN]– consists of 1 C-atom, 1 N-atom, and -1 charge which accounts for 1 extra valence electron. Thus, the valence electrons present in [CN]– Lewis structure = 4 + 5 + 1 = 10 valence electrons.

2. Select the central atom

In the cyanide [CN]– Lewis structure, we do not need to find the least electronegative atom to choose the central atom as we do so in other molecules or molecular ions. As there are only two atoms present in [CN]– so we place them next to each other, as shown below.

3. Connect the two atoms using a single straight line

 In this step, we need to connect the carbon and nitrogen atom using a single straight line, as shown below.

The straight line represents a single covalent bond i.e., a bond pair containing 2 electrons. 2 electrons consumed out of the 10 initially available leaves behind 10-2 = 8 valence electrons. So, we still need to accommodate these 8 electrons in the Lewis dot structure of CN–. Let’s see how we can do so in the upcoming steps.

4. Complete the octet of the more electronegative atom

Nitrogen is more electronegative than carbon. A nitrogen (N) atom needs a total of 8 valence electrons in order to achieve a stable octet electronic configuration. A C-N single bond represents that this N-atom already has 2 electrons.

It is thus short of 6 more electrons to complete its octet. Thus, these 6 electrons are placed as 3 lone pairs around the N-atom, as shown below.

5. Complete the octet of the less electronegative atom

• Total valence electrons used till step 4 = 1 single bond + electrons placed around the N-atom (shown as dots) = 2 + 6 = 8 valence electrons.
• Total valence electrons available – electrons used till step 4 = 10 – 8 = 2.

So these two valence electrons are now placed as a lone pair on the less electronegative carbon atom.

But the problem here is that in the above structure, the C-atom has 1 single bond and 1 lone pair which makes a total of 2 +2 = 4 valence electrons. This means it is still short of 4 valence electrons in order to achieve a complete octet.

This problem can be solved if we convert 2 lone pairs present on the N-atom into two covalent chemical bonds between the bonded C and N-atoms, as shown below.

The figure above shows that both the C and N-atoms now have a complete octet having 1 triple bond + 1 lone pair each.

As a final step, we need to check the stability of this Lewis structure. Let’s do that using the formal charge concept.

6. Check the stability of [CN]–Lewis structure using the formal charge concept

The less the formal charge on the atoms of a molecule or molecular ion, the better the stability of its Lewis structure.

The formal charge can be calculated using the formula given below.

• Formal charge = [ valence electrons – nonbonding electrons- ½ (bonding electrons)]

Let us demonstrate how we can use this formula and the Lewis structure obtained in step 5 to calculate the required formal charges.

For carbon atom

• Valence electrons of carbon = 4
• Bonding electrons = 1 triple bond = 3(2) = 6 electrons
• Non-bonding electrons = 1 lone pair = 2 electrons
• Formal charge = 4-2-6/2 = 4-2-3 = 4-5 = -1

For nitrogen atom

• Valence electrons of nitrogen = 5
• Bonding electrons = 1 triple bond = 3(2) = 6 electrons
• Non-bonding electrons = 1 lone pair = 2 electrons
• Formal charge = 5-2-6/2 = 5-2-3 = 5-5 = 0

The above calculation shows that a zero formal charge is present on the nitrogen atom. Contrarily, a -1 formal charge is present on the carbon atom which is also the charge present on the ion overall. Consequently, the Lewis structure obtained above is correct. It is finally enclosed in square brackets and a -1 charge is placed at the top right corner.

By following the same guidelines as given for drawing [CN]– anion Lewis structure, you can also draw the Lewis structures of the CN molecule and CN+ cation, using a total of 9 and 8 valence electrons respectively.

Now that we have successfully drawn the Lewis dot structure of CN–, let us proceed forward and find out how we can draw its molecular orbital diagram.

Also check –

• How to draw a lewis structure?
• Formal charge calculator
• Lewis structure generator