COF2 Lewis structure

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COF₂ (carbonyl fluoride) has one carbon atom, one oxygen atom, and two fluorine atoms.

In the COF₂ Lewis structure, there is one double bond and two single bonds around the carbon atom, with one oxygen atom and two fluorine atoms attached to it. Each fluorine atom has three lone pairs, and the oxygen atom has two lone pairs.

Steps

To properly draw the COF₂ Lewis structure, follow these steps:

#1 Draw a rough sketch of the structure
#2 Next, indicate lone pairs on the atoms
#3 Indicate formal charges on the atoms, if necessary
#4 Minimize formal charges by converting lone pairs of the atoms
#5 Repeat step 4 if necessary, until all charges are minimized

Let’s break down each step in more detail.

#1 Draw a rough sketch of the structure

• First, determine the total number of valence electrons

In the periodic table, carbon lies in group 14, oxygen lies in group 16, and fluorine lies in group 17.

Hence, carbon has four valence electrons, oxygen has six valence electrons, and fluorine has seven valence electrons.

Since COF₂ has one carbon atom, one oxygen atom, and two fluorine atoms, so…

Valence electrons of one carbon atom = 4 × 1 = 4
Valence electrons of one oxygen atom = 6 × 1 = 6
Valence electrons of two fluorine atoms = 7 × 2 = 14

And the total valence electrons = 4 + 6 + 14 = 24

Learn how to find: Carbon valence electrons, Oxygen valence electrons, and Fluorine valence electrons

• Second, find the total electron pairs

We have a total of 24 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 = 24 ÷ 2 = 12

• Third, determine the central atom

We have to place the least electronegative atom at the center.

Since carbon is less electronegative than oxygen and fluorine, assume that the central atom is carbon.

Therefore, place carbon in the center and oxygen and fluorine on either side.

• And finally, draw the rough sketch

#2 Next, indicate lone pairs on the atoms

Here, we have a total of 8 electron pairs. And two bonds are already marked. So we have to only mark the remaining six electron pairs as lone pairs on the sketch.

Also remember that all three (carbon, oxygen and fluorine) 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 oxygen and fluorines.

So for oxygen and each fluorine, there are three lone pairs, and for carbon, there is zero lone pair because all six electron pairs are over.

Mark the lone pairs on the sketch as follows:

#3 Indicate formal charges on the atoms, if necessary

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

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

For carbon atom, formal charge = 4 – 0 – ½ (6) = +1

For oxygen atom, formal charge = 6 – 6 – ½ (2) = -1

For each fluorine atom, formal charge = 7 – 6 – ½ (2) = 0

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

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

#4 Minimize formal charges by converting lone pairs of the atoms

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

In the above structure, you can see that the central atom (carbon) forms an octet. And the outside atoms (oxygen and fluorines) also form 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 COF₂.

Next: NO₂F Lewis structure

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External links

• https://techiescientist.com/cof2-lewis-structure/
• https://lambdageeks.com/cof2-lewis-structure/
• https://www.thegeoexchange.org/chemistry/bonding/Lewis-Structures/COF2-lewis-structure.html
• https://oneclass.com/homework-help/chemistry/6910298-cof2-molecular-geometry.en.html