HCN Lewis Structure

HCN Lewis Structure

HCN (hydrogen cyanide) has one hydrogen atom, one carbon atom, and one nitrogen atom. In the lewis structure of HCN, there is a single bond between carbon and hydrogen atom, and a triple bond between carbon and nitrogen atom, and on nitrogen atom, there is one lone pair.


Here’s how you can draw the HCN lewis structure step by step.

Step #1: draw sketch
Step #2: mark lone pairs
Step #3: mark charges
Step #4: minimize charges
Step #5: minimize charges again (if there are)

Let’s break down each step in detail.

#1 Draw Sketch

  • First, determine the total number of valence electrons

In the periodic table, hydrogen lies in group 1, carbon lies in group 14, and nitrogen lies in group 15.

Hence, hydrogen has one valence electron, carbon has four valence electrons, and nitrogen has five valence electrons.

Since HCN has one hydrogen atom, one carbon atom, and one nitrogen atom, so…

Valence electrons of one hydrogen atom = 1 × 1 = 1
Valence electrons of one carbon atom = 4 × 1 = 4
Valence electrons of one nitrogen atom = 5 × 1 = 5

And the total valence electrons = 1 + 4 + 5 = 10

  • Second, find the total electron pairs

We have a total of 10 valence electrons. And when we divide this value by two, we get the value of total electron pairs.

Total electron pairs = total valence electrons ÷ 2

So the total electron pairs = 10 ÷ 2 = 5

  • Third, determine the central atom

Here hydrogen can not be the central atom. Because the central atom is bonded with at least two other atoms, and hydrogen has only one electron in its last shell, so it can not make more than one bond.

Now we have to choose the central atom from carbon and nitrogen. Place the least electronegative atom at the center.

Since carbon is less electronegative than nitrogen, assume that the central atom is carbon.

Therefore, place carbon in the center and hydrogen and nitrogen on either side.

  • And finally, draw the rough sketch
HCN Lewis Structure (Step 1)

#2 Mark Lone Pairs

Here, we have a total of 5 electron pairs. And two bonds are already marked. So we have to only mark the remaining three electron pairs as lone pairs on the sketch.

Also remember that hydrogen is a period 1 element, so it can not keep more than 2 electrons in its last shell. And both (carbon and nitrogen) are the period 2 elements, so they can not keep more than 8 electrons in their last shell.

Always start to mark the lone pairs from outside atoms. Here, the outside atoms are hydrogen and nitrogen. But no need to mark on hydrogen, because hydrogen already has two electrons.

So for nitrogen, there are three lone pairs, and for carbon, there is zero lone pair because all three electron pairs are over.

Mark the lone pairs on the sketch as follows:

HCN Lewis Structure (Step 2)

#3 Mark Charges

Use the following formula to calculate the formal charges on atoms:

Formal charge = valence electrons – nonbonding electrons – ½ bonding electrons

For hydrogen atom, formal charge = 1 – 0 – ½ (2) = 0

For carbon atom, formal charge = 4 – 0 – ½ (4) = +2

For nitrogen atom, formal charge = 5 – 6 – ½ (2) = -2

Here, both carbon and nitrogen atoms have charges, so mark them on the sketch as follows:

HCN Lewis Structure (Step 3)

The above structure is not a stable lewis structure because both carbon and nitrogen atoms have charges. Therefore, reduce the charges (as below) by converting lone pairs to bonds.

#4 Minimize Charges

Convert a lone pair of the nitrogen atom to make a new C — N bond with the carbon atom as follows:

HCN Lewis Structure (Step 4)

#5 Minimize Charges Again

Since there are charges on carbon and nitrogen atoms, again convert a lone pair of the nitrogen atom to make a new C — N bond with the carbon atom as follows:

HCN Lewis Structure (Step 5)

In the above structure, you can see that the central atom (carbon) forms an octet. Hence, the octet rule is satisfied.

Also, the above structure is more stable than the previous structures. Therefore, this structure is the stable lewis structure of HCN.

Next: H2O Lewis Structure

Leave a Comment