Question

Which of the following is not a disproportionation reaction?

• A. \(\mathrm{Hg_2Cl_2 \longrightarrow Hg+HgCl_2}\)
• B. \(\mathrm{4H_3PO_3 \longrightarrow 3H_3PO_4+PH_3}\)
• C. \(\mathrm{Cl_2+2NaOH \longrightarrow NaCl+NaOCl+H_2O}\)
• D. \(\mathrm{BaO_2+H_2SO_4 \longrightarrow BaSO_4+H_2O_2}\)

Hint:

To identify disproportionation reactions, track the oxidation number of the same element on both reactant and product sides.

Answer 1

The Correct Option is D

Step 1: Define the criterion for disproportionation.
Disproportionation requires a single element in a single reactant to be simultaneously oxidized (to a higher state) and reduced (to a lower state) in the same reaction.
Step 2: Analyze BaO₂ + H₂SO₄ reaction.
BaO₂ + H₂SO₄ $\rightarrow$ BaSO₄ + H₂O₂. Assign oxidation states: In BaO₂, Ba = +2, O = -1 (peroxide). In H₂SO₄, H = +1, S = +6, O = -2. In H₂O₂, H = +1, O = -1. In BaSO₄, Ba = +2, S = +6, O = -2.
Step 3: Check each element for oxidation-state change.
Ba: +2 in BaO₂ and +2 in BaSO₄ - no change. O in BaO₂: -1 in BaO₂ and -1 in H₂O₂ - no change. H: +1 throughout. S: +6 throughout.
Step 4: Conclude whether redox occurs.
No element changes its oxidation state in this reaction. It is simply a double displacement (metathesis) reaction: the peroxide ion from BaO₂ is transferred to form H₂O₂, and SO₄^2- forms BaSO₄.
Step 5: Confirm it is not disproportionation.
Disproportionation requires one element to both increase and decrease in oxidation state. Here, no element undergoes any change in oxidation state at all. This reaction is neither redox nor disproportionation.
Step 6: Final answer.
BaO₂ + H₂SO₄ $\rightarrow$ BaSO₄ + H₂O₂ is NOT a disproportionation reaction because no oxidation-state changes occur.
\[ \boxed{\text{Option 4: } BaO_2 + H_2SO_4 \rightarrow BaSO_4 + H_2O_2 \text{ is NOT disproportionation}} \]