Questions: For each chemical reaction listed in the table below, decide whether the highlighted atom is being oxidized or reduced. reaction highlighted atom is being... --- --- --- --- oxidized reduced neither oxidized nor reduced H2 S(aq) + 2 NaOH(aq) → Na2 S(aq) + 2 H2 O(l) O O O FeO(s) + CO(g) → Fe(s) + CO2(g) O ○ ○ NH3(aq) + 2 O2(g) → HNO3(aq) + H2 O(l) O ○ O 2 NH3(aq) + H2 SO4(aq) → (NH4)2 SO4(aq) O 0 O

For each chemical reaction listed in the table below, decide whether the highlighted atom is being oxidized or reduced. reaction highlighted atom is being... --- --- --- --- oxidized reduced neither oxidized nor reduced H2 S(aq) + 2 NaOH(aq) → Na2 S(aq) + 2 H2 O(l) O O O FeO(s) + CO(g) → Fe(s) + CO2(g) O ○ ○ NH3(aq) + 2 O2(g) → HNO3(aq) + H2 O(l) O ○ O 2 NH3(aq) + H2 SO4(aq) → (NH4)2 SO4(aq) O 0 O
Transcript text: For each chemical reaction listed in the table below, decide whether the highlighted atom is being oxidized or reduced. \begin{tabular}{|c|c|c|c|} \hline \multirow[b]{2}{*}{reaction} & \multicolumn{3}{|l|}{highlighted atom is being...} \\ \hline & oxidized & reduced & neither oxidized nor reduced \\ \hline \[ \mathrm{H}_{2} \mathrm{~S}(a q)+2 \mathrm{NaOH}(a q) \rightarrow \mathrm{Na}_{2} \mathrm{~S}(a q)+2 \mathrm{H}_{2} \mathrm{O}(I) \] & O & O & O \\ \hline \[ \mathrm{FeO}(\mathrm{~s})+\mathrm{CO}(\mathrm{~g}) \rightarrow \mathrm{Fe}(\mathrm{~s})+\mathrm{CO}_{2}(\mathrm{~g}) \] & O & ○ & ○ \\ \hline \[ \mathrm{NH}_{3}(\mathrm{aq})+2 \mathrm{O}_{2}(\mathrm{~g}) \rightarrow \mathrm{HNO}_{3}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \] & O & ○ & O \\ \hline \[ 2 \mathrm{NH}_{3}(a q)+\mathrm{H}_{2} \mathrm{SO}_{4}(a q) \rightarrow\left(\mathrm{NH}_{4}\right)_{2} \mathrm{SO}_{4}(a q) \] & O & 0 & O \\ \hline \end{tabular}
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Solution

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Solution Steps

Step 1: Understanding Oxidation and Reduction

Oxidation involves the loss of electrons, resulting in an increase in oxidation state, while reduction involves the gain of electrons, resulting in a decrease in oxidation state. To determine if an atom is oxidized or reduced, we need to compare the oxidation states of the highlighted atom before and after the reaction.

Step 2: Analyzing Each Reaction

  1. Reaction 1: \[ \mathrm{H}_{2} \mathrm{~S}(a q)+2 \mathrm{NaOH}(a q) \rightarrow \mathrm{Na}_{2} \mathrm{~S}(a q)+2 \mathrm{H}_{2} \mathrm{O}(I) \]

    • The highlighted atom is sulfur (S).
    • In \(\mathrm{H}_{2}\mathrm{S}\), the oxidation state of S is \(-2\).
    • In \(\mathrm{Na}_{2}\mathrm{S}\), the oxidation state of S is also \(-2\).
    • Since the oxidation state of sulfur does not change, it is neither oxidized nor reduced.

  2. Reaction 2: \[ \mathrm{FeO}(\mathrm{~s})+\mathrm{CO}(\mathrm{~g}) \rightarrow \mathrm{Fe}(\mathrm{~s})+\mathrm{CO}_{2}(\mathrm{~g}) \]

    • The highlighted atom is iron (Fe).
    • In \(\mathrm{FeO}\), the oxidation state of Fe is \(+2\).
    • In \(\mathrm{Fe}\), the oxidation state of Fe is \(0\).
    • The oxidation state of Fe decreases, indicating that Fe is reduced.

  3. Reaction 3: \[ \mathrm{NH}_{3}(\mathrm{aq})+2 \mathrm{O}_{2}(\mathrm{~g}) \rightarrow \mathrm{HNO}_{3}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \]

    • The highlighted atom is nitrogen (N).
    • In \(\mathrm{NH}_{3}\), the oxidation state of N is \(-3\).
    • In \(\mathrm{HNO}_{3}\), the oxidation state of N is \(+5\).
    • The oxidation state of N increases, indicating that N is oxidized.

Final Answer

  1. \(\mathrm{H}_{2} \mathrm{~S}(a q)+2 \mathrm{NaOH}(a q) \rightarrow \mathrm{Na}_{2} \mathrm{~S}(a q)+2 \mathrm{H}_{2} \mathrm{O}(I)\): \(\boxed{\text{neither oxidized nor reduced}}\)
  2. \(\mathrm{FeO}(\mathrm{~s})+\mathrm{CO}(\mathrm{~g}) \rightarrow \mathrm{Fe}(\mathrm{~s})+\mathrm{CO}_{2}(\mathrm{~g})\): \(\boxed{\text{reduced}}\)
  3. \(\mathrm{NH}_{3}(\mathrm{aq})+2 \mathrm{O}_{2}(\mathrm{~g}) \rightarrow \mathrm{HNO}_{3}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l})\): \(\boxed{\text{oxidized}}\)
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