To calculate the number of lone pairs in a molecule like sulfur hexafluoride (SF₆), you can follow these steps:
-
Determine the total number of valence electrons:
- Sulfur (S) is in group 16 and has 6 valence electrons.
- Each fluorine (F) atom is in group 17 and has 7 valence electrons. Since there are 6 fluorine atoms, this contributes \(6 \times 7 = 42\) electrons.
- Total valence electrons for SF₆ = \(6 + 42 = 48\) valence electrons.
-
Draw the Lewis structure:
- Sulfur is the central atom, and it will form single bonds with each of the 6 fluorine atoms.
- Each single bond uses 2 electrons, so for 6 bonds, \(6 \times 2 = 12\) electrons are used.
- Subtract the bonded electrons from the total: \(48 - 12 = 36\) electrons remain.
-
Distribute the remaining electrons:
- The remaining 36 electrons are used to complete the octets of the fluorine atoms. Each fluorine atom needs 6 more electrons to complete its octet (since each already has 2 from the bond with sulfur).
- With 6 fluorine atoms needing 6 electrons each, that uses \(6 \times 6 = 36\) electrons, which uses up all the remaining electrons.
-
Count lone pairs:
- In the Lewis structure of SF₆, sulfur has no lone pairs because all of its valence electrons are used in bonding with fluorine.
- Each fluorine has 3 lone pairs (each having 6 electrons around them).
In summary, in SF₆, the central sulfur atom has 0 lone pairs, while each of the 6 fluorine atoms has 3 lone pairs.
1 answer