# CS2 Lewis structure

CS2 (carbon disulfide) has one carbon atom and two sulfur atoms.

In the CS2 Lewis structure, there are two double bonds around the carbon atom, with two sulfur atoms attached to it, and on each sulfur atom, there are two lone pairs.

Contents

## Steps

Use these steps to correctly draw the CS2 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, carbon lies in group 14, and sulfur lies in group 16.

Hence, carbon has four valence electrons and sulfur has six valence electrons.

Since CS2 has one carbon atom and two sulfur atoms, so…

Valence electrons of one carbon atom = 4 × 1 = 4
Valence electrons of two sulfur atoms = 6 × 2 = 12

And the total valence electrons = 4 + 12 = 16

• Second, find the total electron pairs

We have a total of 16 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 = 16 ÷ 2 = 8

• Third, determine the central atom

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

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

Therefore, place sulfur in the center and sulfurs on either side.

• And finally, draw the rough sketch

### #2 Mark lone pairs on the atoms

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

Also remember that carbon is a period 2 element, so it can not keep more than 8 electrons in its 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 sulfurs.

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

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

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

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

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

### #5 Repeating step 4 to get a stable Lewis structure

Since there are charges on carbon and sulfur atoms, again convert a lone pair of the sulfur atom to make a new C — S 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 (sulfurs) 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 CS2.

Next: SF6 Lewis structure