# SO2F2 Lewis structure

SO2F2 (sulfuryl fluoride) has one sulfur atom, two oxygen atoms, and two fluorine atoms.

In SO2F2 Lewis structure, there are two double bonds and two single bonds around the sulfur atom, with two oxygen atoms and two fluorine atoms attached to it. Each oxygen atom has two lone pairs, and each fluorine atom has three lone pairs.

Contents

## Steps

To properly draw the SO2F2 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, both sulfur and oxygen lie in group 16, and fluorine lies in group 17.

Hence, both sulfur and oxygen have six valence electrons, and fluorine has seven valence electrons.

Since SO2F2 has one sulfur atom, two oxygen atoms, and two fluorine atoms, so…

Valence electrons of one sulfur atom = 6 × 1 = 6
Valence electrons of two oxygen atoms = 6 × 2 = 12
Valence electrons of two fluorine atoms = 7 × 2 = 14

And the total valence electrons = 6 + 12 + 14 = 32

• Second, find the total electron pairs

We have a total of 32 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 = 32 ÷ 2 = 16

• Third, determine the central atom

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

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

Therefore, place sulfur 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 16 electron pairs. And four bonds are already marked. So we have to only mark the remaining twelve electron pairs as lone pairs on the sketch.

Also remember that sulfur is a period 3 element, so it can keep more than 8 electrons in its last shell. And both (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 oxygens and fluorines.

So for each oxygen and each fluorine, there are three lone pairs, and for sulfur, there is zero lone pair because all twelve 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 sulfur atom, formal charge = 6 – 0 – ½ (8) = +2

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

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

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

The above structure is not a stable Lewis structure because both sulfur 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 S — O bond with the sulfur atom as follows:

### #5 Repeat step 4 (minimize charges again)

Since there are charges on sulfur and oxygen atoms, again convert a lone pair of the oxygen atom to make a new S — O bond with the sulfur atom as follows:

In the above structure, you can see that the central atom (sulfur) forms an octet. And the outside atoms (oxygens 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 SO2F2.