Question

Match the following columns.

Column I(Substance)		Column II(Product after electrolysis)
A	Aqueous solution of $AgNO _{3}$ using $Ag$ electrodes	p	Oxygen is produced at anode
B	Aqueous solution of $AgNO _{3}$ using $Pt$ electrodes	q	Hydrogen is produced at cathode
C	Dilute solution of $H _{2} SO _{4}$ at using $Pt$ electrodes	r	Silver is deposited cathode
D	Aqueous solution of $CuCl _{2}$ is using $Pt$ electrodes	s	Neither $O _{2}$ nor $H _{2}$ produced

• A. $A-(r, q), B-(p, q), C-(q, s), D-(p, s)$
• B. $A-(r, q), B-(p, q, s), C-(q, s), D-(q, r, s)$
• C. $A-(p), B-(q), C-(r), D-(s)$
• D. $A-(r, s), B-(r, p), C-(p, q), D-(s)$

Answer 1

Answer: (D)

$A \rightarrow r , s ; B \rightarrow r , p ; C \rightarrow p , q ; D \rightarrow s$
A. An aqueous solution of $AgNO _{3}$ with silver electrodes. In aqueous solution, ionisation of $AgNO _{3}$ and $H _{2} O$ takes place.
$AgNO_{3}(s) \xrightleftharpoons {(aq)}Ag ^{+}(a q)+ NO _{3}^{-}(a q)$
$H _{2} O (l) \rightleftharpoons H ^{+}(a q)+ OH ^{-}(a q)$
At cathode $Ag ^{+}$ions has less discharge potential than $H ^{+}$ ions so silver will be deposited at cathode.
$Ag ^{+}(a q)+e^{-} \longrightarrow Ag(s)$
At anode An equivalent amount of silver will be oxidised to $Ag ^{+}$ions by releasing electrons.
$Ag (s) \longrightarrow Ag ^{+}(a q)+e^{-}$
Ag anode is attacked by $NO _{3}^{-}$ions, so it will also produce $Ag ^{+}$in the solution.
B. An aqueous solution of $AgNO _{3}$ with platinum electrodes. In aqueous solution, ionisation of $AgNO _{3}$ and $H _{2} O$ both occur.
$AgNO_{3}(s) \xrightleftharpoons {(aq)} Ag ^{+}(a q)+ NO _{3}^{-}(a q)$
$H _{2} O (l) \rightleftharpoons H ^{+}(a q)+ OH ^{-}(a q)$
As platinum electrodes are non-attackable electrodes, they will not be reacted upon by $NO _{3}^{-}$ions.
At cathode Ag will be deposited at cathode.
$Ag ^{+}(a q)+e^{-} \longrightarrow Ag(s)$
At anode Out of $NO _{3}^{-}$and $OH ^{-}$ions, only $OH ^{-}$ions will be oxidised (due to less discharge potential) preferentially and $NO _{3}^{-}$ions will remain in the solution.
$OH ^{-}(a q) \longrightarrow OH +e^{-}$
$4 OH \longrightarrow 2 H _{2} O (l)+ O _{2}(g)$
So, oxygen gas is produced at anode. The solution remains acidic due to the presence of $HNO _{3}$.
$H ^{+}(a q)+ NO _{3}^{-}(a q) \rightleftharpoons HNO _{3}(a q)$
C. A dilute solution of $H _{2} SO _{4}$ with platinum electrodes. Both $H _{2} SO _{4}$ and water ionise in the solution.
$H _{2} SO _{4}(a q) \rightleftharpoons 2 H ^{+}(a q)+\operatorname{SO}_{4}^{2-}(a q) $
$H _{2} O (l) \rightleftharpoons 2 H ^{+}(a q)+ OH ^{-}(a q)$
At cathode $H ^{+}$ions will be reduced and hydrogen gas is produced at cathode.
$H ^{+}(a q)+e^{-} \longrightarrow H (g)$
$H (g)+ H (g) \longrightarrow H _{2}(g)$
At anode $OH ^{-}$ions will be released preferentially and not $SO _{4}^{-}$ions due to less discharge potential.
$OH ^{-}(a q) \longrightarrow OH +e^{-}$
$4 OH \longrightarrow 2 H _{2} O (l)+ O _{2}(g)$
Oxygen gas is produced at anode.
Solution will be acidic and will contain $H _{2} SO _{4}$.
D. An aqueous solution of $CuCl _{2}$ with platinum electrodes. Both $CuCl _{2}$ and water ionise as usual.
$CuCl _{2} \xrightleftharpoons {(aq)}Cu ^{2+}(a q)+2 Cl ^{-}(a q)$
$H _{2} O (l) \rightleftharpoons 2 H ^{+}(a q)+ OH ^{-}(a q)$
At cathode $Cu ^{2+}$ ions will be reduced preferentially due to less discharge potential than $H ^{+}$ions.
$Cu ^{2+}(a q)+2 e^{-} \longrightarrow Cu (s)$
Copper metal is deposited at cathode.
At anode $Cl ^{-}$ions will be discharged in preference to $OH ^{-}$ions and chlorine gas is produced at anode.
$Cl ^{-}(a q) \longrightarrow Cl (g)+e^{-}$
$Cl (g)+ Cl (g) \longrightarrow Cl _{2}(g)$