# SOF2 Lewis structure

SOF2 (thionyl fluoride) has one sulfur atom, one oxygen atom, and two fluorine atoms.

In SOF2 Lewis structure, there are two single bonds and one double bond around the sulfur atom, with two fluorine atoms and one oxygen atom attached to it. Each fluorine atom has three lone pairs, the oxygen atom has two lone pairs, and the sulfur atom has one lone pair.

Contents

## Steps

Use these steps to correctly draw the SOF2 Lewis structure:

#1 First draw a rough sketch
#2 Mark lone pairs on the atoms
#3 Calculate and mark formal charges on the atoms, if required
#4 Convert lone pairs of the atoms, and minimize formal charges
#5 Repeat step 4 if needed, until all charges are minimized, to get a stable Lewis structure

Let’s discuss each step in more detail.

### #1 First draw a rough sketch

• 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 SOF2 has one sulfur atom, one oxygen atom, and two fluorine atoms, so…

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

And the total valence electrons = 6 + 6 + 14 = 26

• Second, find the total electron pairs

We have a total of 26 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 = 26 ÷ 2 = 13

• 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 Mark lone pairs on the atoms

Here, we have a total of 13 electron pairs. And three bonds are already marked. So we have to only mark the remaining ten 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 oxygen and fluorines.

So for oxygen and each fluorine, there are three lone pairs, and for sulfur, there is one lone pair.

Mark the lone pairs on the sketch as follows:

### #3 Calculate and mark formal charges on the atoms, if required

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 – 2 – 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 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 Convert lone pairs of the atoms, and minimize formal charges

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 (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 SOF2.