Advertisements
Advertisements
Question
Balance the following ionic equations.
\[\ce{Cr2O^{2-}7 + H^{+} + I- -> Cr^{3+} + I2 + H2O}\]
Solution
Dividing the equation into two half-reactions:
Oxidation half-reaction: \[\ce{I- -> I2}\]
Reduction half-reaction: \[\ce{CrO7^{2-} -> Cr^{3+}}\]
Balancing oxidation and reduction half-reactions separately as:
Oxidation half-reaction:
\[\ce{I- -> I2}\]
\[\ce{2I- -> I2}\]
\[\ce{2I- -> I2 + 2e-}\] ......(i)
Reduction half-reaction:
\[\ce{CrO^{2-}7 -> Cr^{3+}}\]
\[\ce{Cr2O^{2-}7 -> 2Cr^{3+}}\]
\[\ce{Cr2O^{2-}7 + 6e- -> 2Cr^{3+}}\]
\[\ce{Cr2O^{2-}7 + 14H^{+} + 6e- -> \underset{(acidic medium)}{2Cr^{3+} + 7H2O}}\] .....(ii)
To balance the electrons, multiply equation (i) by 3 and add to equation (ii)
\[\ce{Cr2O^{2-}7 + 14H^{+} + 6I- -> 2Cr^{3+} + 3I2 + 7H2O}\]
APPEARS IN
RELATED QUESTIONS
Consider the reaction:
\[\ce{O3(g) + H2O2(l) → H2O(l) + 2O2(g)}\]
Why it is more appropriate to write these reaction as:
\[\ce{O3(g) + H2O2 (l) → H2O(l) + O2(g) + O2(g)}\]
Also, suggest a technique to investigate the path of the redox reactions.
Balance the following redox reactions by ion-electron method:
- \[\ce{MnO-_4 (aq) + I– (aq) → MnO2 (s) + I2(s) (in basic medium)}\]
- \[\ce{MnO-_4 (aq) + SO2 (g) → Mn^{2+} (aq) + HSO-_4 (aq) (in acidic solution)}\]
- \[\ce{H2O2 (aq) + Fe^{2+} (aq) → Fe^{3+} (aq) + H2O (l) (in acidic solution)}\]
- \[\ce{Cr_2O^{2-}_7 + SO2(g) → Cr^{3+} (aq) + SO^{2-}_4 (aq) (in acidic solution)}\]
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.
Justify that the following reaction is redox reaction; identify the species oxidized/reduced, which acts as an oxidant and which acts as a reductant.
\[\ce{2Cu2O_{(S)} + Cu2S_{(S)}->6Cu_{(S)} + SO2_{(g)}}\]
Balance the following reaction by oxidation number method.
\[\ce{H2SO4_{(aq)} + C_{(s)}->CO2_{(g)} + SO2_{(g)} + H2O_{(l)}(acidic)}\]
Balance the following reaction by oxidation number method.
\[\ce{Bi(OH)_{3(s)} + Sn(OH)^-_{3(aq)}->Bi_{(s)} + Sn(OH)^2-_{6(aq)}(basic)}\]
Balance the following redox equation by half-reaction method.
\[\ce{Bi(OH)_{3(s)} + SnO^2-_{2(aq)}->SnO^2-_{3(aq)} + Bi^_{(s)}(basic)}\]
When methane is burnt completely, oxidation state of carbon changes from ______.
Identify the redox reactions out of the following reactions and identify the oxidising and reducing agents in them.
\[\ce{Fe2O3 (s) + 3CO (g) ->[Δ] 2Fe (s) + 3CO2 (g)}\]
In the reaction of oxalate with permanganate in an acidic medium, the number of electrons involved in producing one molecule of CO2 is ______.
