Advertisements
Advertisements
Question
Identify the redox reactions out of the following reactions and identify the oxidising and reducing agents in them.
\[\ce{Fe2O3 (s) + 3CO (g) ->[Δ] 2Fe (s) + 3CO2 (g)}\]
Advertisements
Solution
\[\ce{F\overset{+3}{e2}\overset{-2}{O3} (s) + 3\overset{+2}{C}\overset{-2}{O (g)} ->[Δ] \overset{0}{2F}e (s) + \overset{+4}{3C}\overset{-2}{O2} (g)}\]
Here, O.N. of \[\ce{Fe}\] decreases from +3 (in \[\ce{Fe2O3}\]) to 0 (in \[\ce{Fe}\]) and therefore, \[\ce{Fe2O3}\] acts as an oxidizing agent.
O.N. of \[\ce{C}\] increases from +2 (in \[\ce{CO}\]) to +4 (in \[\ce{CO2}\]) and therefore, \[\ce{CO}\] acts as a reducing agent. Thus, this is a redox reaction.
APPEARS IN
RELATED QUESTIONS
Calculate the oxidation number of sulphur, chromium and nitrogen in H2SO5, `"Cr"_2"O"_7^(2-)` and `"NO"_3^-`. Suggest structure of these compounds. Count for the fallacy.
Whenever a reaction between an oxidising agent and a reducing agent is carried out, a compound of lower oxidation state is formed if the reducing agent is in excess and a compound of higher oxidation state is formed if the oxidising agent is in excess. Justify this statement giving three illustrations.
The Mn3+ ion is unstable in solution and undergoes disproportionation to give Mn2+, MnO2, and H+ ion. Write a balanced ionic equation for the reaction.
Choose the correct option.
For the following redox reactions, find the correct statement.
\[\ce{Sn^{2⊕} + 2Fe^{3⊕}->Sn^{4⊕} + 2Fe^{2⊕}}\]
Balance the following redox equation by half-reaction method.
\[\ce{H2C2O_{4(aq)} + MnO^-_{4(aq)}->CO2_{(g)} + Mn^2+_{( aq)}(acidic)}\]
What is the change in oxidation number of Sulphur in following reaction?
\[\ce{MnO^-_{4(aq)} + SO^{2-}_{3(aq)} -> MnO^{2-}_{4(aq)} + SO^{2-}_{4(aq)}}\]
When methane is burnt completely, oxidation state of carbon changes from ______.
Consider the reaction:
\[\ce{6 CO2(g) + 6H2O(l) → C6 H12O6(aq) + 6O2(g)}\]
Why it is more appropriate to write these reaction as:
\[\ce{6CO2(g) + 12H2O(l) → C6 H12O6(aq) + 6H2O(l) + 6O2(g)}\]
Also, suggest a technique to investigate the path of the redox reactions.
Balance the following equations by the oxidation number method.
\[\ce{Fe^{2+} + H^{+} + Cr2O^{2-}7 -> Cr^{3+} + Fe^{3+} + H2O}\]
Balance the following equations by the oxidation number method.
\[\ce{I2 + NO^{-}3 -> NO2 + IO^{-}3}\]
Balance the following equations by the oxidation number method.
\[\ce{MnO2 + C2O^{2-}4 -> Mn^{2+} + CO2}\]
Identify the redox reactions out of the following reactions and identify the oxidising and reducing agents in them.
\[\ce{HgCl2 (aq) + 2KI (aq) -> HgI2 (s) + 2KCl (aq)}\]
Balance the following ionic equations.
\[\ce{MnO^{-}4 + H^{+} + Br^{-} -> Mn^{2+} + Br2 + H2O}\]
In \[\ce{Cu^{2+} + Ag -> Cu + Ag^+}\], oxidation half-reaction is:
In the reaction of oxalate with permanganate in an acidic medium, the number of electrons involved in producing one molecule of CO2 is ______.
\[\ce{H2O2 -> 2H^+ + O2 + 2e^-}\]; E0 = −0.68 V.
This equation represents which of the following behaviour of H2O2?
On balancing the given redox reaction,
\[\ce{aCr2O7^2- + Bso3^2- (aq) + CH+ (aq) -> 2aCr^3+ (aq) + bSO4^2- (aq) + c(H2O(1)/2}\]
the coefficients a, b and c are found to be, respectively:
