Advertisements
Advertisements
प्रश्न
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.
Advertisements
उत्तर
The given reaction can be represented as:
\[\ce{Mn^{(3+)}_{(aq)} -> Mn^{(2+)}_{(aq)} + MnO_{2(s)} + H+_{(aq)}}\]
The oxidation half-equation is:
\[\ce{^{+3}Mn^{3+}_{(aq)} -> ^{+4}MnO_{2(s)}}\]
The oxidation number is balanced by adding one electron as:
\[\ce{Mn^{3+}_{(aq)} -> MnO_{2(s)} + e-}\]
The charge is balanced by adding 4H+ ions as:
\[\ce{Mn^{3+}_{(aq)} -> MnO_{2(s)} + 4H_{(aq)}^ + e-}\]
The O atoms and H+ ions are balanced by adding 2H2O molecules as:
\[\ce{Mn^{3+}_{(aq)} + 2H2O_{(l)} -> MnO_{2(s)} + 4H+_{(aq)} + e-}\] .....(i)
The reduction half equation is:
\[\ce{Mn^{3+}_{(aq)} -> Mn^{2+}_{(aq)}}\]
The oxidation number is balanced by adding one electron as:
\[\ce{Mn^{3+}_{(aq)} + e- -> Mn^{2+}_{(aq)}}\] ....(ii)
The balanced chemical equation can be obtained by adding equation (i) and (ii) as:
\[\ce{2Mn^{3+}_{(aq)} + 2H_2O_{(l)} -> MnO_{2(s)} + 2Mn^{2+}_{(aq)} + 4H+_{(aq)}}\]
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.
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.
Balance the following redox reactions by ion-electron method:
- \[\ce{MnO-_4 (aq) + I– (aq) → MnO2 (s) + I2(s) (in basic medium)}\]
- \[\ce{MnO-_4 (aq) + SO2 (g) → Mn^{2+} (aq) + HSO-_4 (aq) (in acidic solution)}\]
- \[\ce{H2O2 (aq) + Fe^{2+} (aq) → Fe^{3+} (aq) + H2O (l) (in acidic solution)}\]
- \[\ce{Cr_2O^{2-}_7 + SO2(g) → Cr^{3+} (aq) + SO^{2-}_4 (aq) (in acidic solution)}\]
Choose the correct option.
For the following redox reactions, find the correct statement.
\[\ce{Sn^{2⊕} + 2Fe^{3⊕}->Sn^{4⊕} + 2Fe^{2⊕}}\]
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.
\[\ce{Bi(OH)_{3(s)} + Sn(OH)^-_{3(aq)}->Bi_{(s)} + Sn(OH)^2-_{6(aq)}(basic)}\]
Which of the following is a redox reaction?
Identify the oxidising agent in the following reaction:
\[\ce{CH4_{(g)} + 2O2_{(g)} -> CO2_{(g)} + 2H2O_{(l)}}\]
When methane is burnt completely, oxidation state of carbon changes from ______.
Write balanced chemical equation for the following reactions:
Dichlorine heptaoxide \[\ce{(Cl2O7)}\] in gaseous state combines with an aqueous solution of hydrogen peroxide in acidic medium to give chlorite ion \[\ce{(ClO^{-}2)}\] and oxygen gas. (Balance by ion-electron method)
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.
\[\ce{HgCl2 (aq) + 2KI (aq) -> HgI2 (s) + 2KCl (aq)}\]
Balance the following ionic equations.
\[\ce{Cr2O^{2-}7 + Fe^{2+} + H+ -> Cr^{3+} + Fe^{3+} + H2O}\]
In \[\ce{Cu^{2+} + Ag -> Cu + Ag^+}\], oxidation half-reaction 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:
