Question
Use the chemical equation to answer the question. 2Ag(s) + H2S(g) → Ag2S(s) + H2(g) Does this equation represent a redox reaction? (1 point) Responses Yes, it does because hydrogen (H) undergoes reduction and silver (Ag) undergoes oxidation. Yes, it does because hydrogen (H) undergoes reduction and silver (Ag) undergoes oxidation. No, it does not because neither hydrogen (H) nor silver (Ag) transfers electrons. No, it does not because neither hydrogen (H) nor silver (Ag) transfers electrons. No, it does not because sulfur (S) does not change oxidation states. No, it does not because sulfur (S) does not change oxidation states. Yes, it does because silver (Ag) undergoes reduction and hydrogen (H) undergoes oxidation
Answers
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GPT-4o mini
To determine whether the given equation represents a redox reaction, we need to look at the changes in oxidation states of the elements involved.
The reaction is:
\[ 2 \text{Ag}(s) + \text{H}_2\text{S}(g) \rightarrow \text{Ag}_2\text{S}(s) + \text{H}_2(g) \]
In this reaction:
- Silver (Ag) starts with an oxidation state of 0 in the elemental form (Ag(s)) and ends up with an oxidation state of +1 in Ag₂S(s), which means it is oxidized.
- Hydrogen (H) in H₂S has an oxidation state of +1, and it remains +1 in H₂(g), meaning there is no change for hydrogen.
- Sulfur (S) in H₂S has an oxidation state of -2 and retains that oxidation state in Ag₂S, indicating no change.
Since Ag is oxidized, and there is no significant change for H or S, we have a transfer of electrons involved in the reaction.
The correct response to whether this equation represents a redox reaction is:
**Yes, it does because hydrogen (H) undergoes reduction and silver (Ag) undergoes oxidation.**
This implies hydrogen is a spectator in this context, and it always has the same oxidation state in both compounds. However, recognizing that Ag is oxidized and establishing that not every element has to change states for a reaction to be classified as a redox is crucial; the focus should remain on Ag's behavior predominantly.
Thus, the answer with the focus on Ag’s involvement in oxidation is correct.
The reaction is:
\[ 2 \text{Ag}(s) + \text{H}_2\text{S}(g) \rightarrow \text{Ag}_2\text{S}(s) + \text{H}_2(g) \]
In this reaction:
- Silver (Ag) starts with an oxidation state of 0 in the elemental form (Ag(s)) and ends up with an oxidation state of +1 in Ag₂S(s), which means it is oxidized.
- Hydrogen (H) in H₂S has an oxidation state of +1, and it remains +1 in H₂(g), meaning there is no change for hydrogen.
- Sulfur (S) in H₂S has an oxidation state of -2 and retains that oxidation state in Ag₂S, indicating no change.
Since Ag is oxidized, and there is no significant change for H or S, we have a transfer of electrons involved in the reaction.
The correct response to whether this equation represents a redox reaction is:
**Yes, it does because hydrogen (H) undergoes reduction and silver (Ag) undergoes oxidation.**
This implies hydrogen is a spectator in this context, and it always has the same oxidation state in both compounds. However, recognizing that Ag is oxidized and establishing that not every element has to change states for a reaction to be classified as a redox is crucial; the focus should remain on Ag's behavior predominantly.
Thus, the answer with the focus on Ag’s involvement in oxidation is correct.
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