2 H2O2(aq) → 2 H2O(l) + O2(g) ΔH° = −196 kJ/mol_rxn The decomposition of H2O2(aq) is represented by — Electrochemistry Chemistry Question

STEPS:

1. Recall the rules for assigning oxidation numbers:
* The oxidation number of an element in its free elemental form is always $0$.
* Hydrogen bonded to nonmetals is assigned an oxidation number of $+1$.
* The sum of the oxidation numbers of all atoms in a neutral compound must equal $0$.
2. Assign oxidation numbers to the atoms in the reactant molecule ($\text{H}_2\text{O}_2$):
* In hydrogen peroxide ($\text{H}_2\text{O}_2$), hydrogen is assigned its standard oxidation state of $+1$.
* To maintain electrical neutrality for the compound, the sum of the oxidation states must be zero:
$$2(+1) + 2(\text{Oxygen}) = 0 \implies 2(\text{Oxygen}) = -2 \implies \mathbf{\text{Oxygen} = -1}$$
* Oxygen carries an unusual $-1$ oxidation state in peroxides.
3. Assign oxidation numbers to the atoms in the product molecules ($\text{H}_2\text{O}$ and $\text{O}_2$):
* In water ($\text{H}_2\text{O}$), hydrogen is $+1$ and oxygen has its standard oxidation state of $-2$.
* In diatomic oxygen gas ($\text{O}_2$), oxygen is in its elemental state, meaning its oxidation number is $0$.
4. Compare the reactant and product states to identify redox changes:
* Hydrogen: Starts at $+1$ in $\text{H}_2\text{O}_2$ and ends at $+1$ in $\text{H}_2\text{O}$. Because there is no change in its oxidation state, hydrogen is neither oxidized nor reduced.
* Oxygen: Starts at $-1$ in $\text{H}_2\text{O}_2$:
* Some oxygen atoms are converted into $\text{H}_2\text{O}$, changing from $-1$ to $-2$. This decrease in oxidation number (gain of electrons) represents reduction.
* Some oxygen atoms are converted into $\text{O}_2$, changing from $-1$ to $0$. This increase in oxidation number (loss of electrons) represents oxidation.
5. Conclude the reaction type: Because oxygen is both oxidized and reduced in the same reaction (a process known as a *disproportionation* reaction), Option C is the correct answer.

*

WHY_OTHERS_WRONG:

• Option A is incorrect: This option claims hydrogen is oxidized (which would require its oxidation state to increase above $+1$) and oxygen is reduced. However, hydrogen's oxidation state remains unchanged at $+1$ throughout the process.
• Option B is incorrect: This option claims hydrogen is reduced (which would require its oxidation state to decrease below $+1$). Hydrogen does not change oxidation state in this reaction.
• Option D is incorrect: This option claims that no elements are oxidized or reduced. This is incorrect because the oxidation state of oxygen changes from $-1$ to both $-2$ and $0$, making this a redox reaction.