To determine the correct formula for iron (III) sulfate, we need to understand the oxidation states of the elements involved and how they combine to form the compound.
1. Iron (III) Ion:
- Iron in the (III) state means it has an oxidation number of +3.
- Symbolically, it's noted as \( Fe^{3+} \).
2. Sulfate Ion:
- The sulfate ion is a polyatomic ion with the formula \( SO_4^{2-} \).
- It has a charge of 2-.
To form a neutral compound, the total positive charge from the iron ions must balance the total negative charge from the sulfate ions.
3. Balancing the Charges:
- Each \( Fe^{3+} \) ion has a charge of +3.
- Each \( SO_4^{2-} \) ion has a charge of -2.
- Let’s assume we need \( x \) iron ions and \( y \) sulfate ions to balance the charges:
- \( 3x \) (total positive charge)
- \( 2y \) (total negative charge)
4. Setting Up the Equation:
- To balance the charges: \( 3x = 2y \)
5. Finding the Smallest Whole Numbers:
- Solving \( 3x = 2y \):
- \( x = 2 \) and \( y = 3 \)
So, to balance the charges, we need 2 iron (III) ions (each \( Fe^{3+} \)) and 3 sulfate ions (each \( SO_4^{2-} \)).
Therefore, the formula for iron (III) sulfate is:
[tex]\[ Fe_2(SO_4)_3 \][/tex]
The correct answer is:
C. [tex]\( Fe_2(SO_4)_3 \)[/tex]
