Step 1: Understanding the Concept:
A redox (reduction-oxidation) reaction involves the transfer of electrons from one reactant to another.
This transfer manifests as a change in the oxidation states (or oxidation numbers) of the atoms involved.
If the oxidation states of any elements change from the reactant side to the product side, it is a redox reaction.
Step 2: Key Formula or Approach:
1. Assign oxidation numbers to every element in the reactants and products for each equation.
2. Compare the oxidation states of elements before and after the reaction.
3. If an element's oxidation number increases (loses electrons), it is oxidized.
4. If an element's oxidation number decreases (gains electrons), it is reduced.
5. A reaction must have both oxidation and reduction to be a redox reaction.
Step 3: Detailed Explanation:
Let's check each option by assigning oxidation states:
Option (A): \( \text{NaCl} + \text{KNO}_3 \rightarrow \text{NaNO}_3 + \text{KCl} \)
Reactants: \( \text{Na} = +1, \text{Cl} = -1, \text{K} = +1, \text{N} = +5, \text{O} = -2 \)
Products: \( \text{Na} = +1, \text{N} = +5, \text{O} = -2, \text{K} = +1, \text{Cl} = -1 \)
No oxidation states change. This is a double displacement reaction.
Option (B): \( \text{Mg(OH)}_2 + 2\text{NH}_4\text{Cl} \rightarrow \text{MgCl}_2 + 2\text{NH}_4\text{OH} \)
Reactants: \( \text{Mg} = +2, \text{O} = -2, \text{H} = +1, \text{N} = -3, \text{Cl} = -1 \)
Products: \( \text{Mg} = +2, \text{Cl} = -1, \text{N} = -3, \text{H} = +1, \text{O} = -2 \)
No oxidation states change. This is an acid-base/double displacement reaction.
Option (C): \( \text{CaC}_2\text{O}_4 + 2\text{HCl} \rightarrow \text{CaCl}_2 + \text{H}_2\text{C}_2\text{O}_4 \)
Reactants: \( \text{Ca} = +2, \text{C} = +3, \text{O} = -2, \text{H} = +1, \text{Cl} = -1 \)
Products: \( \text{Ca} = +2, \text{Cl} = -1, \text{H} = +1, \text{C} = +3, \text{O} = -2 \)
No oxidation states change. This is a double displacement reaction.
Option (D): \( \text{Zn} + 2\text{AgCN} \rightarrow 2\text{Ag} + \text{Zn(CN)}_2 \)
Reactants:
- \( \text{Zn} \) is a free, uncombined element, so its oxidation state is 0.
- In \( \text{AgCN} \), \( \text{Ag} \) is +1, and the cyanide ion (\( \text{CN}^- \)) as a whole is -1.
Products:
- \( \text{Ag} \) is now a free, uncombined element, so its oxidation state is 0.
- In \( \text{Zn(CN)}_2 \), \( \text{Zn} \) is +2 (to balance the two -1 cyanide ions).
Changes:
- Zinc (\( \text{Zn} \)) goes from 0 to +2. Its oxidation number increased, so it was oxidized.
- Silver (\( \text{Ag} \)) goes from +1 to 0. Its oxidation number decreased, so it was reduced.
Because oxidation numbers changed, this is a redox reaction (specifically, a single displacement).
Step 4: Final Answer:
The reaction \( \text{Zn} + 2\text{AgCN} \rightarrow 2\text{Ag} + \text{Zn(CN)}_2 \) is a redox reaction.