Question

What will be the oxidation number of the elements in \(O_3\), \(P_4\) and \(S_8\)?

• A. \(-1,0,+1\)
• B. \(1,+1,-2\)
• C. \(0,0,0\)
• D. \(-2,1,0\)

Hint:

The oxidation number of any element in its free or elemental state is always zero, regardless of whether it exists as monoatomic, diatomic, or polyatomic molecules.

Answer 1

The Correct Option is C

Concept: The oxidation number of an element represents the apparent charge that an atom would possess if all bonds were considered completely ionic. For elements present in their free or elemental state, the oxidation number is always zero irrespective of the molecular complexity of the species. Examples of elemental forms include: \[ H_2,\quad O_2,\quad O_3,\quad P_4,\quad S_8,\quad N_2 \] Even though these molecules contain multiple atoms bonded together, all atoms belong to the same element and share electrons equally. Therefore, no atom acquires a positive or negative oxidation state.

Step 1: Determine the oxidation number of oxygen in ozone \((O_3)\). Ozone is an allotrope of oxygen consisting only of oxygen atoms. \[ O_3 \] Since it is the elemental form of oxygen, every oxygen atom has oxidation number: \[ 0 \]

Step 2: Determine the oxidation number of phosphorus in \(P_4\). White phosphorus exists as: \[ P_4 \] Since phosphorus is present in its elemental form, the oxidation number of each phosphorus atom is: \[ 0 \]

Step 3: Determine the oxidation number of sulfur in \(S_8\). Sulfur commonly exists as: \[ S_8 \] This is also an elemental form of sulfur. Therefore, \[ \text{Oxidation number of sulfur}=0 \] Hence, \[ O_3 : 0,\qquad P_4 : 0,\qquad S_8 : 0 \] Therefore, the correct option is \[ \boxed{(c)} \]