Formal charge is calculated using the formula: Formal Charge = (Valence electrons) - (Non-bonding electrons) - (1/2)(Bonding electrons). In the most basic Lewis structure of SO_3, sulfur forms double bonds with two oxygen atoms and a single bond with the third. Sulfur has 6 valence electrons. If sulfur forms one double bond and two single bonds (with complete octets on all oxygens), sulfur has 0 non-bonding electrons and shares 8 bonding electrons (2 double bond + 2 single bond pairs = not quite). Let us be precise: in a structure where S has two double bonds and one single bond, S uses 10 electrons in bonding (five pairs) with no lone pairs. Formal charge on S = 6 - 0 - (10/2) = 6 - 5 = +1. However, in the simplest structure with all single bonds and a positive charge, or with resonance involving S=O bonds, the formal charge on S varies. In the Lewis structure with three S=O double bonds (expanded octet), formal charge = 6 - 0 - 6 = 0, but this requires 12 electrons around S. In the most commonly drawn single-resonance structure with one double bond and two coordinate/dative bonds from S to O (S^+ to O^-), formal charge on S = 6 - 0 - (8/2) = +2. This is the textbook answer for the un-expanded octet structure. Option +1 applies when S forms two double bonds. Option 0 applies in the expanded-octet structure. Option -1 is not a correct formal charge for sulfur in SO_3. This is a standard NCERT/JEE concept on formal charge and resonance. Plausibility: SO_3 is a known electrophile, consistent with positive formal charge on S in the basic Lewis structure.