Advertisements
Advertisements
प्रश्न
Balance the following ionic equations.
\[\ce{MnO^{-}4 + SO^{2-}3 + H^{+} -> Mn^{2+} + SO^{2-}4 + H2O}\]
Advertisements
उत्तर
Dividing the equation into two half-reactions:
Oxidation half-reaction: \[\ce{SO^{2-}3 -> SO^{2-}4}\]
Reduction half-reaction: \[\ce{MnO^{-}4 -> MN^{2+}}\]
Balancing oxidation and reduction half-reactions separately as:
Oxidation half-reaction:
\[\ce{SO^{2-}3 -> SO^{2-}4}\]
\[\ce{SO^{2-}3 -> SO^{2-}4 + 2e-}\]
Since the reaction occurs in acidic medium,
\[\ce{SO^{2-}3 -> SO^{2-}4 + 2e^{-} + 2H+}\]
\[\ce{SO^{2-}3 + H2O -> SO^{2-}4 + 2H^{+} + 2e-}\] .....(i)
Reduction half-reaction:
\[\ce{MnO^{-}4 -> MN^{2+}}\]
\[\ce{MnO^{-}4 + 5e^{-} -> MN^{2+}}\]
\[\ce{MnO^{-}4 + 5e^{-} -> MN^{2+} + 4H2O}\] .....(ii)
To balance the electrons, multiply equation (i) by 5 and equation (ii) by 2 and add
\[\ce{2MO^{-}4 + 5SO^{2-}3 + 6H+ -> 2Mn^{2+} + 5SO^{2-}4 + 3H2O}\]
APPEARS IN
संबंधित प्रश्न
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.
The compound AgF2 is an unstable compound. However, if formed, the compound acts as a very strong oxidizing agent. Why?
Balance the following equation in basic medium by ion-electron method and oxidation number methods and identify the oxidising agent and the reducing agent.
\[\ce{P4(s) + OH–(aq) —> PH3(g) + HPO^–_2(aq)}\]
Balance the following equation in basic medium by ion-electron method and oxidation number methods and identify the oxidising agent and the reducing agent.
\[\ce{Cl_2O_{7(g)} + H_2O_{2(aq)} -> ClO-_{2(aq)} + O_{2(g)} + H+_{(aq)}}\]
Balance the following reaction by oxidation number method.
\[\ce{H2SO4_{(aq)} + C_{(s)} -> CO2_{(g)} + SO2_{(g)} + H2O_{(l)}(acidic)}\]
Balance the following redox equation by half-reaction method.
\[\ce{H2C2O_{4(aq)} + MnO^-_{4(aq)}->CO2_{(g)} + Mn^2+_{( aq)}(acidic)}\]
Balance the following redox equation by half-reaction method.
\[\ce{Bi(OH)_{3(s)} + SnO^2-_{2(aq)}->SnO^2-_{3(aq)} + Bi^_{(s)}(basic)}\]
Balance the following equations by the oxidation number method.
\[\ce{I2 + S2O^{2-}3 -> I- + S4O^{2-}6}\]
Identify the redox reactions out of the following reactions and identify the oxidising and reducing agents in them.
\[\ce{HgCl2 (aq) + 2KI (aq) -> HgI2 (s) + 2KCl (aq)}\]
Identify the redox reactions out of the following reactions and identify the oxidising and reducing agents in them.
\[\ce{Fe2O3 (s) + 3CO (g) ->[Δ] 2Fe (s) + 3CO2 (g)}\]
Identify the redox reactions out of the following reactions and identify the oxidising and reducing agents in them.
\[\ce{4NH3 (g) + 3O2 (g) -> 2N2 (g) + 6H2O (g)}\]
Balance the following ionic equations.
\[\ce{Cr2O^{2-}7 + Fe^{2+} + H+ -> Cr^{3+} + Fe^{3+} + H2O}\]
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:
The weight of CO is required to form Re2(CO)10 will be ______ g, from 2.50 g of Re2O7 according to given reaction
\[\ce{Re2O7 + CO -> Re2(CO)10 + CO2}\]
Atomic weight of Re = 186.2; C = 12 and O = 16.
\[\ce{H2O2 -> 2H^+ + O2 + 2e^-}\]; E0 = −0.68 V.
This equation represents which of the following behaviour of H2O2?
