XeF3– has one xenon atom and three fluorine atoms. In the lewis structure of XeF3–, there are three single bonds around the xenon atom, with three fluorine atoms attached to it, and each atom has three lone pairs.
Also, there is a negative (-1) charge on the xenon atom.
Here’s how you can draw the XeF3– 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
Hence, xenon has eight valence electrons and fluorine has seven valence electrons.
Since XeF3– has one xenon atom and three fluorine atoms, so…
Valence electrons of one xenon atom = 8 × 1 = 8
Valence electrons of three fluorine atoms = 7 × 3 = 21
Now the XeF3– has a negative (-1) charge, so we have to add one more electron.
So the total valence electrons = 8 + 21 + 1 = 30
- Second, find the total electron pairs
We have a total of 30 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 = 30 ÷ 2 = 15
- Third, determine the central atom
We have to place the least electronegative atom at the center.
Since xenon is less electronegative than fluorine, assume that the central atom is xenon.
Therefore, place xenon in the center and fluorines on either side.
- And finally, draw the rough sketch
#2 Mark Lone Pairs
Here, we have a total of 15 electron pairs. And three Xe — F bonds are already marked. So we have to only mark the remaining twelve electron pairs as lone pairs on the sketch.
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 xenon, there are three lone pairs.
Mark the lone pairs on the sketch as follows:
#3 Mark Charges
Use the following formula to calculate the formal charges on atoms:
Formal charge = valence electrons – nonbonding electrons – ½ bonding electrons
For xenon atom, formal charge = 8 – 6 – ½ (6) = -1
For each fluorine atom, formal charge = 7 – 6 – ½ (2) = 0
Here, the xenon atom has a charge, so mark it on the sketch as follows:
In the above structure, you can see that the central atom (xenon) forms an octet. Hence, the octet rule is satisfied.
Now there is still a negative (-1) charge on the xenon 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 xenon atom to make a new Xe — 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 XeF3–.
And since the XeF3– has a negative (-1) charge, mention that charge on the lewis structure by drawing brackets as follows:
Next: XeF5+ Lewis Structure