Lewis structure of PCl4+

PCl₄⁺ has one phosphorus atom and four chlorine atoms.

In the Lewis structure of PCl₄⁺, there are four single bonds around the phosphorus atom, with four chlorine atoms attached to it, and on each chlorine atom, there are three lone pairs.

Also, there is a positive (+1) charge on the phosphorus atom.

Alternative method: PCl₄⁺ Lewis structure

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 positive (+1) charge, so we have to subtract one electron.

So the total valence electrons = 5 + 28 – 1 = 32

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

• Second, find the total electron pairs

We have a total of 32 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 = 32 ÷ 2 = 16

• 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 16 electron pairs. And four P — Cl bonds are already marked. So we have to only mark the remaining twelve 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 zero lone pair because all twelve electron pairs are over.

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 – 0 – 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⁺ has a central phosphorus atom connected to four chlorine atoms through single covalent bonds. Within this layout, the phosphorus atom satisfies the octet rule by forming four bonding pairs with no lone pairs. Each chlorine atom fulfills its octet by maintaining three lone pairs of its own alongside the single shared bond. This arrangement 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. Therefore, this specific electronic distribution serves as the definitive and most accurate Lewis representation of the PCl₄⁺ ion.

To properly represent this as a polyatomic ion, the entire Lewis structure is enclosed within square brackets. The overall charge of 1+ is then written as a superscript outside the brackets at the top right, indicating that the structure possesses one fewer electron than the total valence count of the five neutral atoms.

Next: BrO^– Lewis structure

External video

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

External links

• https://lambdageeks.com/pcl4-lewis-structure/
• https://www.bartleby.com/questions-and-answers/whats-the-lewis-dot-structure-for-pcl-4-ion/ee8fd50c-a0ac-4277-9717-4c30cd1b251c
• https://www.numerade.com/ask/question/whats-the-lewis-dot-structure-for-pcl4-ion-70628/