PF6- Lewis Structure

PF6- Lewis Structure

PF6 (hexafluorophosphate) has one phosphorus atom and six fluorine atoms. In the lewis structure of PF6, there are six single bonds around the phosphorus atom, with six fluorine atoms attached to it, and on each fluorine atom, there are three lone pairs.

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

Steps

Here’s how you can draw the PF6 lewis structure step by step.

Step #1: draw sketch
Step #2: mark lone pairs
Step #3: mark charges (if there are)

Let’s break down each step in detail.

#1 Draw Sketch

  • 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 PF6 has one phosphorus atom and six fluorine atoms, so…

Valence electrons of one phosphorus atom = 5 × 1 = 5
Valence electrons of six fluorine atoms = 7 × 6 = 42

Now the PF6 has a negative (-1) charge, so we have to add one more electron.

So the total valence electrons = 5 + 42 + 1 = 48

  • Second, find the total electron pairs

We have a total of 48 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 = 48 ÷ 2 = 24

  • 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
PF6- Lewis Structure (Step 1)

#2 Mark Lone Pairs

Here, we have a total of 24 electron pairs. And six P — F bonds are already marked. So we have to only mark the remaining eighteen 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 eighteen electron pairs are over.

Mark the lone pairs on the sketch as follows:

PF6- Lewis Structure (Step 2)

#3 Mark Charges

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 (12) = -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:

PF6- Lewis Structure (Step 3)

In the above structure, you can see that the central atom (phosphorus) forms 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 fluorine.

But if we convert a lone pair of the phosphorus atom to make a new P — F bond with the fluorine 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 PF6.

And since the PF6 has a negative (-1) charge, mention that charge on the lewis structure by drawing brackets as follows:

PF6- Lewis Structure (Final)

Next: H3PO4 Lewis Structure

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