PF2- Lewis structure

PF₂^– has one phosphorus atom and two fluorine atoms.

In PF₂^– Lewis structure, there are two single bonds around the phosphorus atom, with two fluorine atoms attached to it. Each fluorine atom has three lone pairs, and the phosphorus atom has two lone pairs.

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

Alternative method: Lewis structure of PF₂^–

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 PF₂^– has one phosphorus atom and two fluorine atoms, so…

Valence electrons of one phosphorus atom = 5 × 1 = 5
Valence electrons of two fluorine atoms = 7 × 2 = 14

Now the PF₂^– has a negative (-1) charge, so we have to add one more electron.

So the total valence electrons = 5 + 14 + 1 = 20

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

• Second, find the total electron pairs

We have a total of 20 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 = 20 ÷ 2 = 10

• 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 10 electron pairs. And two P — F bonds are already marked. So we have to only mark the remaining eight 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 are two lone pairs.

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 – 4 – ½ (4) = -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 PF₂^– features a central phosphorus atom connected to two fluorine atoms through single covalent bonds. In this arrangement, the phosphorus atom satisfies the octet rule by forming two bonding pairs and retaining two lone pairs. Each fluorine atom fulfills its octet by maintaining three lone pairs of its own alongside the single shared bond. This configuration is the most stable because it optimizes the formal charge distribution; the phosphorus atom carries a formal charge of -1, while both fluorine atoms maintain a formal charge of zero. Consequently, this specific electronic distribution serves as the definitive and most accurate Lewis representation of the PF₂^– ion.

To complete the representation, draw square brackets around the entire Lewis structure and place a “-” or “-1” sign as a superscript outside the upper right bracket. This notation signifies that the negative charge is a property of the whole ion.

Next: SI₄ Lewis structure

External links

• https://homework.study.com/explanation/draw-the-most-important-lewis-structure-for-pf2-assuming-it-exists-and-then-answer-the-following-questions-the-underlined-atom-is-the-central-atom-all-other-atoms-are-bonded-directly-to-the-c.html
• https://www.chegg.com/homework-help/questions-and-answers/draw-important-lewis-structure-pf2-assuming-exists-answer-following-questions-underlined-a-q17609579