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Question
Justify that the following reaction is redox reaction:
\[\ce{4 NH3(g) + 5 O2(g) → 4NO(g) + 6H2O(g)}\]
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Solution
The oxidation number of each element in the given reaction can be represented as:
-3 +1 0 +2 -2 +1 -2
\[\ce{4 NH3(g) + 5 O2(g) → 4NO(g) + 6H2O(g)}\]
Here, the oxidation number of N increases from - 3 in NH3 to +2 in NO. On the other hand, the oxidation number of O2 decreases from 0 in O2 to - 2 in NO and H2O i.e., O2 is reduced. Hence, the given reaction is a redox reaction.
RELATED QUESTIONS
Justify that the following reaction is redox reaction:
\[\ce{CuO(s) + H2(g) → Cu(s) + H2O(g)}\]
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\[\ce{Fe2O3(s) + 3CO(g) → 2Fe(s) + 3CO2(g)}\]
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When concentrated sulphuric acid is added to an inorganic mixture containing chloride, we get colourless pungent-smelling gas HCl, but if the mixture contains bromide then we get red vapour of bromine. Why?
Identify the substance oxidised, reduced, oxidising agent and reducing agent for the following reaction:
\[\ce{HCHO (l) + 2Cu^{2+}(aq) + 5 OH–(aq) → Cu2O(s) + HCOO–(aq) + 3H2O(l)}\]
Consider the reactions:
\[\ce{2S_2O_3^{(2-)}(aq) + l_2(S) -> S_4O_6^{(2-)}(aq) + 2l-(aq)}\]
\[\ce{S_2O_3^{(2-)}(aq) + 2Br_2(l) + 5H_2O(l) -> 2SO_4^{2-} (aq) + 4Br-(aq) + 10H+ (aq)}\]
Why does the same reductant, thiosulphate react differently with iodine and bromine?
Consider the reactions:
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- \[\ce{H3PO2(aq) + 2CuSO4(aq) + 2 H2O(l) → H3PO4(aq) + 2Cu(s) + H2SO4(aq)}\]
- \[\ce{C6H5CHO(l) + 2[Ag (NH3)2]+(aq) + 3OH–(aq) → C6H5COO–(aq) + 2Ag(s) + 4NH3 (aq) + 2 H2O(l)}\]
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What inference do you draw about the behaviour of Ag+ and Cu2+ from these reactions?
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Identify disproportionation reaction
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\[\ce{2KClO3 -> 2KCl + 3O2}\]
(i) Potassium is undergoing oxidation.
(ii) Chlorine is undergoing oxidation.
(iii) Oxygen is reduced.
(iv) None of the species are undergoing oxidation or reduction.
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\[\ce{P4 + 3OH- + 3H2O -> PH3 + 3H2PO^{-}2}\]
(i) Phosphorus is undergoing reduction only.
(ii) Phosphorus is undergoing oxidation only.
(iii) Phosphorus is undergoing oxidation as well as reduction.
(iv) Hydrogen is undergoing neither oxidation nor reduction.
Assertion (A): The decomposition of hydrogen peroxide to form water and oxygen is an example of disproportionation reaction.
Reason (R): The oxygen of peroxide is in –1 oxidation state and it is converted to zero oxidation state in \[\ce{O2}\] and –2 oxidation state in \[\ce{H2O}\].
Why does fluorine not show disporportionation reaction?
Which of the following statement is CORRECT for the decomposition reaction of KClO3?
\[\ce{2KClO3 → 2KCl +3O2}\]
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\[\ce{H2O2 -> H2O + O2}\]
This represents ______.
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