PF6- Lewis structure

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PF₆^– (hexafluorophosphate) has one phosphorus atom and six fluorine atoms.

In the PF₆^– Lewis structure, 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 easily draw the PF₆^– Lewis structure step by step:

#1 Draw a rough skeleton structure
#2 Mention lone pairs on the atoms
#3 If needed, mention formal charges on the atoms

Now, let’s take a closer look at each step mentioned above.

#1 Draw a rough skeleton 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 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

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

• 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

#2 Mention lone pairs on the atoms

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:

#3 If needed, mention formal charges on the atoms

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:

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.

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:

Next: H₃PO₄ Lewis structure

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External links

• https://www.thegeoexchange.org/chemistry/bonding/Lewis-Structures/PF6-Lewis-structure.html
• https://lambdageeks.com/pf6-lewis-structures/
• https://oneclass.com/homework-help/chemistry/7027553-pf6-lewis-structure.en.html