PCl4- Lewis structure

PCl₄^– has one phosphorus atom and four chlorine atoms.

In PCl₄^– Lewis structure, there are four single bonds around the phosphorus atom, with four chlorine atoms attached to it. Each chlorine atom has three lone pairs, and the phosphorus atom has one lone pair.

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

Alternative method: Lewis structure of PCl₄^–

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 PCl₄^– has one phosphorus atom and four chlorine atoms, so…

Valence electrons of one phosphorus atom = 5 × 1 = 5
Valence electrons of four chlorine atoms = 7 × 4 = 28

Now the PCl₄^– has a negative (-1) charge, so we have to add one more electron.

So the total valence electrons = 5 + 28 + 1 = 34

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

• Second, find the total electron pairs

We have a total of 34 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 = 34 ÷ 2 = 17

• 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

Lone pair

Here, we have a total of 17 electron pairs. And four P — Cl bonds are already marked. So we have to only mark the remaining thirteen 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 is one lone pair.

Mark the lone pairs on the sketch as follows:

Formal charge

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 – 2 – ½ (8) = -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:

Final structure

The final structure of PCl₄^– involves a central phosphorus atom connected to four chlorine atoms through single covalent bonds. In this arrangement, the phosphorus atom serves as an exception to the octet rule, utilizing an expanded valence shell to accommodate ten electrons, which include four bonding pairs and one lone pair. Each chlorine atom fulfills its octet by maintaining three lone pairs of its own alongside the single shared bond. This configuration is the most stable because it optimizes the formal charge distribution; the phosphorus atom carries a formal charge of -1, while all four chlorine atoms maintain a formal charge of zero. Accordingly, this specific electronic distribution serves as the definitive and most accurate Lewis representation of the PCl₄^– ion.

To complete the representation, draw square brackets around the entire Lewis structure and place a “-” or “-1” sign as a superscript outside the upper right bracket. This notation signifies that the negative charge is a property of the whole ion.

Next: Lewis structure of PCl₄⁺

External video

• PCl4- Lewis Structure – How to Draw the Lewis Structure for PCl4 – – YouTube • Wayne Breslyn

External links

• https://www.thegeoexchange.org/chemistry/bonding/Lewis-Structures/PCl4-minus-lewis-structure.html