PF4+ Lewis structure

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PF₄⁺ has one phosphorus atom and four fluorine atoms.

In PF₄⁺ Lewis structure, there are four single bonds around the phosphorus atom, with four fluorine atoms attached to it, and on each fluorine atom, there are three lone pairs.

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

Steps

To properly draw the PF₄⁺ 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

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

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

Since PF₄⁺ has one phosphorus atom and four fluorine atoms, so…

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

Now the PF₄⁺ 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 Fluorine 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 fluorine, assume that the central atom is phosphorus.

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

• And finally, draw the rough sketch

#2 Next, indicate lone pairs on the atoms

Here, we have a total of 16 electron pairs. And four P — F bonds are already marked. So we have to only mark the remaining twelve electron pairs as lone pairs on the sketch.

Also remember that phosphorus is a period 3 element, so it can keep more than 8 electrons in its last shell. And fluorine is a period 2 element, so it can not keep more than 8 electrons in its last shell.

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

So for each fluorine, 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:

#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 phosphorus atom, formal charge = 5 – 0 – ½ (8) = +1

For each fluorine 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 (fluorines) also form an octet. Hence, the octet rule is satisfied.

Now there is still a positive (+1) charge on the phosphorus atom.

This is okay, because the structure with a positive charge on the least electronegative atom is the best Lewis structure. And in this case, the least electronegative element is phosphorus.

Therefore, this structure is the most stable Lewis structure of PF₄⁺.

And since the PF₄⁺ has a positive (+1) charge, mention that charge on the Lewis structure by drawing brackets as follows:

FAQs

Next: Lewis structure of PF₄^–

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