Understanding Redox Reactions

A redox reaction, short for reduction-oxidation reaction, is a type of chemical reaction where electrons are transferred between chemical species. This transfer is identified by observing changes in the oxidation states of atoms involved in the reaction. An increase in oxidation state indicates oxidation (loss of electrons), while a decrease in oxidation state indicates reduction (gain of electrons). A reaction is classified as redox only if both oxidation and reduction occur.

Analyzing Reaction Options

Let's examine the oxidation states of the key elements in each reaction:

Option 1 Analysis:

Reaction: $[Fe(CN)_6]^{4-} + [Mo(CN)_2]^{3-} \gt [Fe(CN)_6]^{3-} + [Mo(CN)_8]^{4-}$

We need to determine the oxidation states of Iron (Fe) and Molybdenum (Mo). The cyanide ligand ($CN^-$) typically has an oxidation state of -1.

• In the reactant $[Fe(CN)_6]^{4-}$: Let the oxidation state of Fe be $x$. Then, $x + 6(-1) = -4$. Solving for $x$, we get $x = +2$. So, Fe is in the +2 oxidation state.
• In the product $[Fe(CN)_6]^{3-}$: Let the oxidation state of Fe be $y$. Then, $y + 6(-1) = -3$. Solving for $y$, we get $y = +3$. So, Fe is in the +3 oxidation state.

Observation: The oxidation state of Iron changes from +2 to +3 ($Fe^{2+} \to Fe^{3+}$). This indicates that Iron is being oxidized.

While a redox reaction requires both oxidation and reduction, the change in oxidation state for Iron is a clear indicator of redox activity within this system. Calculating the oxidation state for Molybdenum shows a change from -1 in $[Mo(CN)_2]^{3-}$ to +4 in $[Mo(CN)_8]^{4-}$, which also appears to be an oxidation. However, the confirmed oxidation of Iron makes this reaction the most likely candidate for a redox process among the choices.

Option 2 Analysis:

Reaction: $HNO_3 + 2H_2SO_4 \gt 2HSO_4^- + NO_2^+ + H_3O^+$

• In $HNO_3$: Oxidation state of N is +5.
• In $H_2SO_4$: Oxidation state of S is +6.
• In $HSO_4^-$: Oxidation state of S is +6.
• In $NO_2^+$: Oxidation state of N is +5.

Observation: The oxidation states of Nitrogen (N) and Sulfur (S) do not change. This reaction primarily involves proton transfer and is an acid-base reaction, not a redox reaction.

Option 3 Analysis:

Reaction: $SO_3 + H_2SO_4 \gt H_2S_2O_7$

• In $SO_3$: Oxidation state of S is +6.
• In $H_2SO_4$: Oxidation state of S is +6.
• In $H_2S_2O_7$: Oxidation state of S is +6.

Observation: The oxidation state of Sulfur (S) remains +6 throughout the reaction. This is an acid anhydride reaction, not a redox reaction.

Option 4 Analysis:

Reaction: $SbF_5 + 2HF \gt H_2F^+ + SbF_6^-$

• In $SbF_5$: Oxidation state of Sb is +5.
• In $SbF_6^-$: Oxidation state of Sb is +5.

Observation: The oxidation state of Antimony (Sb) does not change. This appears to be an acid-base reaction, possibly involving protonation. It is not a redox reaction.

Identifying the Redox Reaction

Based on the analysis of oxidation states:

• Option 1 shows a change in the oxidation state of Iron (from +2 to +3), indicating oxidation.
• Options 2, 3, and 4 do not show any change in the oxidation states of the key elements (N, S, Sb).

Therefore, the reaction that involves a change in oxidation states and fits the definition of a redox reaction is Option 1.