SF5- Lewis structure

SF₅^– has one sulfur atom and five fluorine atoms.

In the SF₅^– Lewis structure, there are five single bonds around the sulfur atom, with five fluorine atoms attached to it. Each fluorine atom has three lone pairs, and the sulfur atom has one lone pair.

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

Alternative method: Lewis structure of SF₅^–

Rough sketch

• First, determine the total number of valence electrons

In the periodic table, sulfur lies in group 16, and fluorine lies in group 17.

Hence, sulfur has six valence electrons and fluorine has seven valence electrons.

Since SF₅^– has one sulfur atom and five fluorine atoms, so…

Valence electrons of one sulfur atom = 6 × 1 = 6
Valence electrons of five fluorine atoms = 7 × 5 = 35

Now the SF₅^– has a negative (-1) charge, so we have to add one more electron.

So the total valence electrons = 6 + 35 + 1 = 42

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

• Second, find the total electron pairs

We have a total of 42 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 = 42 ÷ 2 = 21

• Third, determine the central atom

We have to place the least electronegative atom at the center.

Since sulfur is less electronegative than fluorine, assume that the central atom is sulfur.

Therefore, place sulfur in the center and fluorines on either side.

• And finally, draw the rough sketch

Lone pair

Here, we have a total of 21 electron pairs. And five S — F bonds are already marked. So we have to only mark the remaining sixteen electron pairs as lone pairs on the sketch.

Also remember that sulfur 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 sulfur, there is one lone pair.

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 sulfur atom, formal charge = 6 – 2 – ½ (10) = -1

For each fluorine atom, formal charge = 7 – 6 – ½ (2) = 0

Here, the sulfur atom has a charge, so mark it on the sketch as follows:

Final structure

The final structure of SF₅^– has a central sulfur atom linked to five fluorine atoms through single covalent bonds. In this arrangement, the sulfur atom utilizes an expanded octet to accommodate twelve valence electrons, which includes five bonding pairs and one lone pair. Each fluorine atom satisfies the octet rule by maintaining three lone pairs of its own. This configuration is the most stable because it results in a formal charge of -1 on the sulfur atom and zero on each of the fluorine atoms, representing the most energetically favorable distribution for the species. Consequently, this specific electronic pattern serves as the definitive and most accurate Lewis representation for this anion.

​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: Lewis structure of SF₃⁺

External video

• How to Draw the Lewis Dot Structure for SF5 – – YouTube • Wayne Breslyn

External links

• https://oneclass.com/homework-help/chemistry/6928944-sf5-lewis-dot-structure.en.html