PF6- Lewis structure

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.

Alternative method: Lewis structure of PF₆^–

Rough 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

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

Lone pair

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:

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 (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:

Final structure

The final structure of PF₆^– contains a central phosphorus atom linked to six fluorine atoms through single covalent bonds. In this configuration, the phosphorus atom utilizes an expanded valence shell to accommodate twelve electrons through these six bonding pairs. Within this layout, each of the six fluorine atoms fulfills the octet rule by maintaining three lone pairs alongside its single shared bond. This arrangement represents the most stable state for the ion because it results in a formal charge of -1 on the phosphorus atom, while each fluorine atom carries a formal charge of zero. Thus, this specific electronic distribution serves as the definitive and most accurate Lewis representation of the hexafluorophosphate 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 additional electron beyond the valence count of the neutral atoms.

Next: H₃PO₄ Lewis structure

External video

• PF6 Lewis Structure: How to Draw the Lewis Structure for Hexafluorophosphate – YouTube • Wayne Breslyn

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