PCl2- Lewis structure

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PCl₂^– has one phosphorus atom and two chlorine atoms.

In PCl₂^– Lewis structure, there are two single bonds around the phosphorus atom, with two chlorine atoms attached to it. Each chlorine atom has three lone pairs, and the phosphorus atom has two lone pairs.

Also, there is a negative (-1) charge on the phosphorus atom.

Steps

Use these steps to correctly draw the PCl₂^– 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

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, phosphorus lies in group 15, and chlorine lies in group 17.

Hence, phosphorus has five valence electrons and chlorine has seven valence electrons.

Since PCl₂^– has one phosphorus atom and two chlorine atoms, so…

Valence electrons of one phosphorus atom = 5 × 1 = 5
Valence electrons of two chlorine atoms = 7 × 2 = 14

Now the PCl₂^– has a negative (-1) charge, so we have to add one more electron.

So the total valence electrons = 5 + 14 + 1 = 20

Learn how to find: Phosphorus valence electrons and Chlorine valence electrons

• Second, find the total electron pairs

We have a total of 20 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 = 20 ÷ 2 = 10

• Third, determine the central atom

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

Since phosphorus is less electronegative than chlorine, assume that the central atom is phosphorus.

Therefore, place phosphorus in the center and chlorines on either side.

• And finally, draw the rough sketch

#2 Mark lone pairs on the atoms

Here, we have a total of 10 electron pairs. And two P — Cl bonds are already marked. So we have to only mark the remaining eight electron pairs as lone pairs on the sketch.

Also remember that both (phosphorus and chlorine) are the period 3 elements, so they can keep more than 8 electrons in their last shell.

Always start to mark the lone pairs from outside atoms. Here, the outside atoms are chlorines.

So for each chlorine, there are three lone pairs, and for phosphorus, there are two lone pairs.

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 phosphorus atom, formal charge = 5 – 4 – ½ (4) = -1

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

Here, the phosphorus atom has a charge, so mark it on the sketch as follows:

In the above structure, you can see that the central atom (phosphorus) forms an octet. And the outside atoms (chlorines) also form an octet. Hence, the octet rule is satisfied.

Now there is still a negative (-1) charge on the phosphorus atom.

This is not okay, right? Because the structure with a negative charge on the most electronegative atom is the best Lewis structure. And in this case, the most electronegative element is chlorine.

But if we convert a lone pair of the phosphorus atom to make a new P — Cl bond with the chlorine atom, and calculate the formal charge, then we do not get the formal charges on atoms closer to zero.

And the structure with the formal charges on atoms closer to zero is the best Lewis structure.

Therefore, this structure is the most stable Lewis structure of PCl₂^–.

And since the PCl₂^– has a negative (-1) charge, mention that charge on the Lewis structure by drawing brackets as follows:

Next: AsO₂^– Lewis structure

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