To determine if the given chemical equation represents a redox reaction, we need to analyze the changes in oxidation states of the elements involved:
Chemical equation:
\[ 2Ag(s) + H2S(g) \rightarrow Ag2S(s) + H2(g) \]
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Oxidation states:
- In elemental silver (\(Ag\)), the oxidation state is 0.
- In hydrogen sulfide (\(H2S\)), the oxidation states are:
- \(H\): +1
- \(S\): -2
- In silver sulfide (\(Ag2S\)), the oxidation states are:
- \(Ag\): +1 (since there are 2 Ag, the total charge needs to balance the -2 from S)
- \(S\): -2
- In molecular hydrogen (\(H2\)), the oxidation state is 0.
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Changes:
- Silver (\(Ag\)) changes from 0 in \(Ag\) to +1 in \(Ag2S\) (oxidation).
- Hydrogen changes from +1 in \(H2S\) to 0 in \(H2\) (reduction).
Based on these changes, we can conclude that silver is oxidized and hydrogen is reduced, indicating that the reaction involves the transfer of electrons.
Conclusion: Yes, it does because silver (Ag) undergoes oxidation and hydrogen (H) undergoes reduction.
The correct response would be:
Yes, it does because silver (Ag) undergoes reduction and hydrogen (H) undergoes oxidation. (One of the options which states that silver undergoes reduction seems incorrect; make sure to select the appropriate one, which should relate to oxidation and reduction accurately.)
3 answers