To draw the Lewis structure for the thiol group (SH), we first need to determine the total number of valence electrons available for the molecule. For sulfur (S), which is in Group 16 of the periodic table, it has 6 valence electrons. Hydrogen (H) has 1 valence electron. Therefore, for SH:
- Valence electrons from sulfur: 6
- Valence electrons from hydrogen: 1
Total valence electrons = 6 + 1 = 7.
Next, we start drawing the structure. Place sulfur in the center because it is less electronegative than hydrogen. Draw a single bond between sulfur and hydrogen. This bond uses 2 of the 7 valence electrons, leaving us with 5 electrons.
Now, we need to complete the octet (or duet for hydrogen) for sulfur. Add lone pairs to the sulfur atom. Use 4 of the remaining 5 valence electrons to create two lone pairs around sulfur. This configuration gives sulfur a total of 8 electrons (2 from the bond with hydrogen and 4 from the lone pairs).
The resulting Lewis structure looks like this:
H: S
:
Now, to determine the formal charge, we use the formula:
Formal Charge = Valence Electrons – (Non-bonding Electrons + 1/2 Bonding Electrons)
For sulfur, the calculation is:
- Valence Electrons = 6
- Non-bonding Electrons = 4 (from the two lone pairs)
- Bonding Electrons = 2 (from the single bond with hydrogen, which is counted as 1 for the sulfur)
So, the formal charge on sulfur is:
Formal Charge = 6 – (4 + 1) = 1.
For hydrogen, the calculation is:
- Valence Electrons = 1
- Non-bonding Electrons = 0
- Bonding Electrons = 2 (the bond with sulfur counts as 1 for hydrogen)
So, the formal charge on hydrogen is:
Formal Charge = 1 – (0 + 1) = 0.
In summary, the Lewis structure for SH shows sulfur bonded to hydrogen with a formal charge of +1 on sulfur and a formal charge of 0 on hydrogen.