Lewis structure of PF6-

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The Lewis structure of PF₆^– contains six single bonds, with phosphorus in the center, and six fluorines on either side. There are three lone pairs on each fluorine atom, and the phosphorus atom does not have any lone pair.

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

Steps

By using the following steps, you can easily draw the Lewis structure of PF₆^–.

#1 Draw skeleton
#2 Show chemical bond
#3 Mark lone pairs
#4 Calculate formal charge and check stability (if octet is already completed on central atom)

Let’s one by one discuss each step in detail.

#1 Draw skeleton

In this step, first calculate the total number of valence electrons. And then, decide the central atom.

• Let’s calculate the total number of valence electrons

We know that… phosphorus is a group 15 element and fluorine is a group 17 element. Hence, phosphorus has five valence electrons and fluorine has seven valence electrons.

Now PF₆^– has one phosphorus atom and six fluorine atoms.

So the total number of valence electrons = valence electrons of phosphorus atom + (valence electrons of fluorine atom × 6)

And PF₆^– has a negative (-1) charge, so we have to add one more electron.

Therefore, the total number of valence electrons = 5 + 42 + 1 = 48

• Now decide the central atom

The atom with the least electronegative value is placed at the center. By looking at the periodic table, we get the electronegativity values for phosphorus and fluorine as follows:

Electronegativity value of phosphorus = 2.19
Electronegativity value of fluorine = 3.98

Obviously, phosphorus is less electronegative than fluorine. Hence, assume that phosphorus is the central atom.

So now, put phosphorus in the center and fluorines on either side. And draw the rough skeleton structure for the Lewis structure of PF₆^– something like this:

Also read: How to draw Lewis structure of XeO₄ (5 steps)

#2 Show chemical bond

Place two electrons between the atoms to show a chemical bond. Since phosphorus is surrounded by six fluorines, use twelve electrons to show six chemical bonds as follows:

Also read: How to draw Lewis structure of CO₃^2- (5 steps)

#3 Mark lone pairs

As calculated earlier, we have a total of 48 valence electrons. And in the above structure, we have already used twelve valence electrons. Hence, thirty-six valence electrons are remaining.

Two valence electrons represent one lone pair. So thirty-six valence electrons = eighteen lone pairs.

Note 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.

Also, make sure that you start marking these lone pairs on outside atoms first. And then, on the central atom.

The outside atoms are fluorines, so each fluorine will get three lone pairs. And the central atom (phosphorus) will not get any lone pair, because all eighteen lone pairs are used.

So the Lewis structure of PF₆^– looks something like this:

In the above structure, you can see that the octet is completed on the central atom (phosphorus), and also on the outside atoms. Therefore, the octet rule is satisfied.

After completing the octet, one last thing we need to do is, calculate the formal charge and check the stability of the above structure.

Also read: How to draw Lewis structure of SO₄^2- (5 steps)

#4 Calculate formal charge and check stability

The following formula is used to calculate the formal charges on atoms:

Formal charge = valence electrons – nonbonding electrons – ½ bonding electrons

Collect the data from the above structure and then, write it down below as follows:

• For phosphorus atom

Valence electrons = 5
Nonbonding electrons = 0
Bonding electrons = 12

Formal charge = 5 – 0 – ½ (12) = -1

• For each fluorine atom

Valence electrons = 7
Nonbonding electrons = 6
Bonding electrons = 2

Formal charge = 7 – 6 – ½ (2) = 0

Mention the formal charges of atoms on the structure. So the Lewis structure of PF₆^– looks something like this:

In the above structure, you can see that the formal charges of atoms are closer to zero. Therefore, this is the most stable Lewis structure of PF₆^–.

And each horizontal line drawn in the above structure represents a pair of bonding valence electrons.

Now PF₆^– is an ion having a negative (-1) charge, so draw brackets around the above Lewis structure and mention that charge on the top right corner. And then, the Lewis structure of PF₆^– looks something like this:

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Related

• Lewis structure of XeO₄
• Lewis structure of CO₃^2-
• Lewis structure of SO₄^2-
• Lewis structure of SiS₂
• Lewis structure of PBr₅

External links

• PF6- Lewis Structure – How to draw the Electron Dot Structure for PF6- – The Geoexchange
• PF6- Lewis Structure in 5 Steps (With Images) – Pediabay
• PF6- Lewis Structure & Characteristics: 11 Complete Facts – Lambda Geeks
• VSEPR calculation for hexafluorophosphate, [PF6]- – University of Sheffield
• PF6- Lewis Structure (Hexafluorophosphate Ion) – Pinterest
• Draw the Lewis structure of the phosphorus hexafluoride ion (PF6) that has the least amount of formal charge on the atoms – Chegg
• Draw the Lewis structure for the phosphorus hexafluoride PF6- ion – Numerade
• PF6- (Hexafluorophosphate Ion) Oxidation Number – ChemicalAid
• Draw the lewis structure of PF6-? – OneClass
• Draw the Lewis structure for the phosphorus hexaflouride (PF6) ion – Bartleby
• What is the Lewis structure of PF₆^–? – Quizlet
• Does PF6- have a dipole moment? – Quora