# NSF Lewis structure

NSF (thiazyl fluoride) has one nitrogen atom, one sulfur atom, and one fluorine atom.

In NSF Lewis structure, there is a triple bond between sulfur and nitrogen atom, and a single bond between sulfur and fluorine atom. The fluorine atom has three lone pairs, and the nitrogen atom and sulfur atom has one lone pair.

Contents

## Steps

To properly draw the NSF 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, nitrogen lies in group 15, sulfur lies in group 16, and fluorine lies in group 17.

Hence, nitrogen has five valence electrons, sulfur has six valence electrons, and fluorine has seven valence electrons.

Since NSF has one nitrogen atom, one sulfur atom, and one fluorine atom, so…

Valence electrons of one nitrogen atom = 5 × 1 = 5
Valence electrons of one sulfur atom = 6 × 1 = 6
Valence electrons of one fluorine atom = 7 × 1 = 7

And the total valence electrons = 5 + 6 + 7 = 18

• Second, find the total electron pairs

We have a total of 18 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 = 18 ÷ 2 = 9

• Third, determine the central atom

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

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

Therefore, place sulfur in the center and nitrogen 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 9 electron pairs. And two bonds are already marked. So we have to only mark the remaining seven electron pairs as lone pairs on the sketch.

Also remember that both (nitrogen and fluorine) are the period 2 elements, so they can not keep more than 8 electrons in their last shell. And sulfur is a period 3 element, so it can keep more than 8 electrons in its last shell.

Always start to mark the lone pairs from outside atoms. Here, the outside atoms are nitrogen and fluorine.

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

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

For sulfur atom, formal charge = 6 – 2 – ½ (4) = +2

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

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

The above structure is not a stable Lewis structure because both nitrogen and sulfur 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 nitrogen atom to make a new S — N bond with the sulfur atom as follows:

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

Since there are charges on the atoms, again convert a lone pair of the nitrogen atom to make a new S — N 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 (nitrogen and fluorine) 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 NSF.