Which balanced equation represents a redox reaction (i.e., involves a change in oxidation numbers)?

• A. $H\mathrm{Cl}$(aq) $+$ $\mathrm{Na}OH$(aq) $\to$ $\mathrm{Na}\mathrm{Cl}$(aq) $+$ $H_{2}O$(l)
• B. ${\mathrm{Na}}_{2}C{O}_3$(aq) $+$ ${\mathrm{Ca}}_{\mathrm{Cl}}{O}_2$(aq) $\to$ ${\mathrm{Ca}}_{1}C{O}_3$(s) $+$ $2\mathrm{Na}\mathrm{Cl}$(aq)
• C. $H_2$(g) $+$ ${\mathrm{Cl}}_2$(g) $\to$ $2H\mathrm{Cl}$(g)
• D. ${\mathrm{Zn}}_1$(s) $+$ ${\mathrm{Cu}}_1{S}_O{O}_4$(aq) $\to$ ${\mathrm{Zn}}_1{S}_O{O}_4$(aq) $+$ ${\mathrm{Cu}}_1$(s)

Answer: (D)

Verified step by step guidance

1. Identify the oxidation states of all elements in the reactants and products for each given equation. This involves assigning oxidation numbers based on known rules (e.g., elemental form = 0, alkali metals = +1, oxygen = -2, etc.).
2. Compare the oxidation states of each element before and after the reaction to see if any element undergoes a change in oxidation number. A redox reaction requires at least one element to be oxidized (increase in oxidation number) and another to be reduced (decrease in oxidation number).
3. For the first equation (HCl + NaOH → NaCl + H₂O), check the oxidation states of H, Cl, Na, and O. Since this is an acid-base neutralization, oxidation states typically do not change, so it is not a redox reaction.
4. For the second equation (Na₂CO₃ + CaCl₂ → CaCO₃ + 2 NaCl), analyze the oxidation states of Na, C, O, Ca, and Cl. This is a double displacement reaction with no change in oxidation numbers, so it is not a redox reaction.
5. For the third equation (H₂ + Cl₂ → 2 HCl), check the oxidation states of H and Cl. Both elements change oxidation states (H from 0 to +1, Cl from 0 to -1), indicating a redox reaction. Similarly, for the fourth equation (Zn + CuSO₄ → ZnSO₄ + Cu), Zn changes from 0 to +2 and Cu from +2 to 0, confirming it is a redox reaction.