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Q2.18 (iv):
Predict the products of electrolysis in each of the following:
(iv) An aqueous solution of (CuCl_{2}) with platinum electrodes.
Solution :
Step 1: Identification of Ionic Species in Aqueous Solution
In an aqueous solution of copper(II) chloride ($CuCl_{2}$), the substance dissociates completely into its constituent ions. Furthermore, water ($H_{2}O$) is present as the solvent and may also undergo electrolysis.
The dissociation reaction is:
$CuCl_{2}(aq) \rightarrow Cu^{2+}(aq) + 2Cl^{-}(aq)$
The auto-ionization/presence of water provides:
$H_{2}O(l) \rightleftharpoons H^{+}(aq) + OH^{-}(aq)$
The species present in the electrolytic cell are: $Cu^{2+}$, $Cl^{-}$, $H^{+}$, $OH^{-}$, and $H_{2}O$.
Step 2: Analysis of Cathodic Reduction Reactions
At the cathode (negative electrode), reduction occurs. We compare the standard reduction potentials ($E^\circ$) of the competing cations:
- Reaction A: $Cu^{2+}(aq) + 2e^{-} \rightarrow Cu(s)$ ; $E^\circ = +0.34\,V$
- Reaction B: $2H_{2}O(l) + 2e^{-} \rightarrow H_{2}(g) + 2OH^{-}(aq)$ ; $E^\circ = -0.83\,V$
[Per the principle of selective discharge, the species with the higher standard reduction potential is preferentially reduced]. Since $E^\circ(Cu^{2+}/Cu) > E^\circ(H_{2}O/H_{2})$, copper ions are preferentially reduced at the cathode.
Step 3: Analysis of Anodic Oxidation Reactions
At the anode (positive electrode), oxidation occurs. We compare the standard oxidation potentials of the competing anions:
- Reaction A: $2Cl^{-}(aq) \rightarrow Cl_{2}(g) + 2e^{-}$ ; $E^\circ_{ox} = -1.36\,V$
- Reaction B: $2H_{2}O(l) \rightarrow O_{2}(g) + 4H^{+}(aq) + 4e^{-}$ ; $E^\circ_{ox} = -1.23\,V$
[Applying the concept of Overpotential]: Although the thermodynamic oxidation potential of water suggests it should be oxidized, the overpotential required for the evolution of oxygen gas makes the oxidation of $Cl^{-}$ kinetically favorable in concentrated solutions.
Step 4: Schematic Representation of Electrolytic Cell
Step 5: Final Summary of Reactions
| Electrode | Reaction | Product |
|---|---|---|
| Cathode | $Cu^{2+}(aq) + 2e^{-} \rightarrow Cu(s)$ | Copper metal ($Cu$) |
| Anode | $2Cl^{-}(aq) \rightarrow Cl_{2}(g) + 2e^{-}$ | Chlorine gas ($Cl_{2}$) |
Final Solution: The electrolysis of aqueous $CuCl_{2}$ using platinum electrodes results in the deposition of metallic copper ($Cu$) at the cathode and the evolution of chlorine gas ($Cl_{2}$) at the anode.
More Questions from Class 12 Chemistry Electrochemistry Exercises
- Q2.1: Arrange the following metals in the order in which they displace each other from the solution of their salts. Al, Cu, Fe, Mg and Zn.
- Q2.1: Arrange the following metals in the order in which they displace each other from the solution of their salts. Al, Cu, Fe, Mg and Zn.
- Q2.10: The conductivity of sodium chloride at 298 K has been determined at different concentrations and the results are given below: Calculate (\Lambda_{m}) for all concentrations and draw a plot between (\Lambda_{m}) and (c^{1/2}). Find the value of (\Lambda^{0}_{m}).
- Q2.10: The conductivity of sodium chloride at 298 K has been determined at different concentrations and the results are given below: Calculate \(\Lambda_{m}\) for all concentrations and draw a plot between \(\Lambda_{m}\) and \(c^{1/2}\). Find the value of \(\Lambda_{m}^{\circ}\).
- Q2.11: Conductivity of 0.00241 M acetic acid is 7.896 (\times) (10^{-5}) S (cm^{-1}). Calculate its molar conductivity. If (\Lambda^{0}_{m}) for acetic acid is 390.5 S (cm^{2} mol^{-1}), what is its dissociation constant?
- Q2.11: Conductivity of 0.00241 M acetic acid is \(7.896 \times 10^{-5} S cm^{-1}\). Calculate its molar conductivity. If \(\Lambda_{m}^{\circ}\) for acetic acid is \(390.5 S cm^{2} mol^{-1}\), what is its dissociation constant?
- Q2.12 (i): How much charge is required for the following reduction: (i) 1 mol of (Al^{3+}) to Al?
- Q2.12 (ii): How much charge is required for the following reduction: (ii) 1 mol of (Cu^{2+}) to Cu?
- Q2.12 (iii): How much charge is required for the following reduction: (iii) 1 mol of (MnO_{4}^{-}) to (Mn^{2+})?
- Q2.12(i): How much charge is required for the following reduction: 1 mol of \(Al^{3+}\) to Al?
- Q2.12(ii): How much charge is required for the following reduction: 1 mol of \(Cu^{2+}\) to Cu?
- Q2.12(iii): How much charge is required for the following reduction: 1 mol of \(MnO_{4}^{-}\) to \(Mn^{2+}\)?
- Q2.13 (i): How much electricity in terms of Faraday is required to produce (i) 20.0 g of Ca from molten (CaCl_{2})?
- Q2.13 (ii): How much electricity in terms of Faraday is required to produce (ii) 40.0 g of Al from molten (Al_{2}O_{3})?
- Q2.13(i): How much electricity in terms of Faraday is required to produce 20.0 g of Ca from molten \(CaCl_{2}\)?
- Q2.13(ii): How much electricity in terms of Faraday is required to produce 40.0 g of Al from molten \(Al_{2}O_{3}\)?
- Q2.14 (i): How much electricity is required in coulomb for the oxidation of (i) 1 mol of (H_{2}O) to (O_{2})?
- Q2.14 (ii): How much electricity is required in coulomb for the oxidation of (ii) 1 mol of (FeO) to (Fe_{2}O_{3})?
- Q2.14(i): How much electricity is required in coulomb for the oxidation of 1 mol of \(H_{2}O\) to \(O_{2}\)?
- Q2.14(ii): How much electricity is required in coulomb for the oxidation of 1 mol of FeO to \(Fe_{2}O_{3}\)?
- Q2.15: A solution of (Ni(NO_{3})_{2}) is electrolysed between platinum electrodes using a current of 5 amperes for 20 minutes. What mass of Ni is deposited at the cathode?
- Q2.15: A solution of \(Ni(NO_{3})_{2}\) is electrolysed between platinum electrodes using a current of 5 amperes for 20 minutes. What mass of Ni is deposited at the cathode?
- Q2.16: Three electrolytic cells A,B,C containing solutions of (ZnSO_{4}), (AgNO_{3}) and (CuSO_{4}), respectively are connected in series. A steady current of 1.5 amperes was passed through them until 1.45 g of silver deposited at the cathode of cell B. How long did the current flow? What mass of copper and zinc were deposited?
- Q2.16: Three electrolytic cells A,B,C containing solutions of \(ZnSO_{4}\), \(AgNO_{3}\) and \(CuSO_{4}\), respectively are connected in series. A steady current of 1.5 amperes was passed through them until 1.45 g of silver deposited at the cathode of cell B. How long did the current flow? What mass of copper and zinc were deposited?
- Q2.17 (i): Using the standard electrode potentials, predict if the reaction between the following is feasible: (i) (Fe^{3+}(aq)) and (I^{-}(aq))
- Q2.17 (ii): Using the standard electrode potentials, predict if the reaction between the following is feasible: (ii) (Ag^{+} (aq)) and Cu(s)
- Q2.17 (iii): Using the standard electrode potentials, predict if the reaction between the following is feasible: (iii) (Fe^{3+} (aq)) and (Br^{-} (aq))
- Q2.17 (iv): Using the standard electrode potentials, predict if the reaction between the following is feasible: (iv) Ag(s) and (Fe^{3+} (aq))
- Q2.17 (v): Using the standard electrode potentials, predict if the reaction between the following is feasible: (v) (Br_{2} (aq)) and (Fe^{2+} (aq)).
- Q2.17(i): Using the standard electrode potentials given in Table 3.1, predict if the reaction between the following is feasible: \(Fe^{3+}(aq)\) and \(I^{-}(aq)\)
- Q2.17(ii): Using the standard electrode potentials given in Table 3.1, predict if the reaction between the following is feasible: \(Ag^{+} (aq)\) and \(Cu(s)\)
- Q2.17(iii): Using the standard electrode potentials given in Table 3.1, predict if the reaction between the following is feasible: \(Fe^{3+} (aq)\) and \(Br^{-} (aq)\)
- Q2.17(iv): Using the standard electrode potentials given in Table 3.1, predict if the reaction between the following is feasible: \(Ag(s)\) and \(Fe^{3+} (aq)\)
- Q2.17(v): Using the standard electrode potentials given in Table 3.1, predict if the reaction between the following is feasible: \(Br_{2} (aq)\) and \(Fe^{2+} (aq)\).
- Q2.18 (i): Predict the products of electrolysis in each of the following: (i) An aqueous solution of (AgNO_{3}) with silver electrodes.
- Q2.18 (ii): Predict the products of electrolysis in each of the following: (ii) An aqueous solution of (AgNO_{3}) with platinum electrodes.
- Q2.18 (iii): Predict the products of electrolysis in each of the following: (iii) A dilute solution of (H_{2}SO_{4}) with platinum electrodes.
- Q2.18(i): Predict the products of electrolysis in each of the following: An aqueous solution of \(AgNO_{3}\) with silver electrodes.
- Q2.18(ii): Predict the products of electrolysis in each of the following: An aqueous solution of \(AgNO_{3}\) with platinum electrodes.
- Q2.18(iii): Predict the products of electrolysis in each of the following: A dilute solution of \(H_{2}SO_{4}\) with platinum electrodes.
- Q2.18(iv): Predict the products of electrolysis in each of the following: An aqueous solution of \(CuCl_{2}\) with platinum electrodes.
- Q2.2: Given the standard electrode potentials, (K^{+}/K) = –2.93V, (Ag^{+}/Ag) = 0.80V, (Hg^{2+}/Hg) = 0.79V (Mg^{2+}/Mg) = –2.37 V, (Cr^{3+}/Cr) = – 0.74V Arrange these metals in their increasing order of reducing power.
- Q2.2: Given the standard electrode potentials, \(K^{+}/K = –2.93V\), \(Ag^{+}/Ag = 0.80V\), \(Hg^{2+}/Hg = 0.79V\) \(Mg^{2+}/Mg = –2.37 V\), \(Cr^{3+}/Cr = – 0.74V\) Arrange these metals in their increasing order of reducing power.
- Q2.3 (i): Depict the galvanic cell in which the reaction (Zn(s)+2Ag^{+}(aq) \rightarrow Zn^{2+}(aq)+2Ag(s)) takes place. Further show: (i) Which of the electrode is negatively charged?
- Q2.3 (ii): Depict the galvanic cell in which the reaction (Zn(s)+2Ag^{+}(aq) \rightarrow Zn^{2+}(aq)+2Ag(s)) takes place. Further show: (ii) The carriers of the current in the cell.
- Q2.3 (iii): Depict the galvanic cell in which the reaction (Zn(s)+2Ag^{+}(aq) \rightarrow Zn^{2+}(aq)+2Ag(s)) takes place. Further show: (iii) Individual reaction at each electrode.
- Q2.3(i): Depict the galvanic cell in which the reaction \(Zn(s)+2Ag^{+}(aq) \rightarrow Zn^{2+}(aq)+2Ag(s)\) takes place. Further show: Which of the electrode is negatively charged?
- Q2.3(ii): Depict the galvanic cell in which the reaction \(Zn(s)+2Ag^{+}(aq) \rightarrow Zn^{2+}(aq)+2Ag(s)\) takes place. Further show: The carriers of the current in the cell.
- Q2.3(iii): Depict the galvanic cell in which the reaction \(Zn(s)+2Ag^{+}(aq) \rightarrow Zn^{2+}(aq)+2Ag(s)\) takes place. Further show: Individual reaction at each electrode.
- Q2.4 (i): Calculate the standard cell potentials of galvanic cell in which the following reaction takes place: (i) (2Cr(s) + 3Cd^{2+}(aq) \rightarrow 2Cr^{3+}(aq) + 3Cd) Calculate the (\Delta_{r}G^{o}) and equilibrium constant of the reaction.
- Q2.4 (ii): Calculate the standard cell potentials of galvanic cell in which the following reaction takes place: (ii) (Fe^{2+}(aq) + Ag^{+}(aq) \rightarrow Fe^{3+}(aq) + Ag(s)) Calculate the (\Delta_{r}G^{o}) and equilibrium constant of the reaction.
- Q2.4(i): Calculate the standard cell potentials of galvanic cell in which the following reaction takes place: \(2Cr(s) + 3Cd^{2+}(aq) \rightarrow 2Cr^{3+}(aq) + 3Cd(s)\) Calculate the \(\Delta_{r}G^{\circ}\) and equilibrium constant of the reaction.
- Q2.4(ii): Calculate the standard cell potentials of galvanic cell in which the following reaction takes place: \(Fe^{2+}(aq) + Ag^{+}(aq) \rightarrow Fe^{3+}(aq) + Ag(s)\) Calculate the \(\Delta_{r}G^{\circ}\) and equilibrium constant of the reaction.
- Q2.5 (i): Write the Nernst equation and emf of the following cell at 298 K: (i) (Mg(s)|Mg^{2+}(0.001M)||Cu^{2+}(0.0001 M)|Cu(s))
- Q2.5 (ii): Write the Nernst equation and emf of the following cell at 298 K: (ii) (Fe(s)|Fe^{2+}(0.001M)||H^{+}(1M)|H_{2}(g)(1bar)| Pt(s))
- Q2.5 (iii): Write the Nernst equation and emf of the following cell at 298 K: (iii) (Sn(s)|Sn^{2+}(0.050 M)||H^{+}(0.020 M)|H_{2}(g) (1 bar)|Pt(s))
- Q2.5 (iv): Write the Nernst equation and emf of the following cell at 298 K: (iv) (Pt(s)|Br^{-}(0.010 M)|Br_{2}(l )||H^{+}(0.030 M)| H_{2}(g) (1 bar)|Pt(s)).
- Q2.5(i): Write the Nernst equation and emf of the following cell at 298 K: \(Mg(s)|Mg^{2+}(0.001M)||Cu^{2+}(0.0001 M)|Cu(s)\)
- Q2.5(ii): Write the Nernst equation and emf of the following cell at 298 K: \(Fe(s)|Fe^{2+}(0.001M)||H^{+}(1M)|H_{2}(g)(1bar)| Pt(s)\)
- Q2.5(iii): Write the Nernst equation and emf of the following cell at 298 K: \(Sn(s)|Sn^{2+}(0.050 M)||H^{+}(0.020 M)|H_{2}(g) (1 bar)|Pt(s)\)
- Q2.5(iv): Write the Nernst equation and emf of the following cell at 298 K: \(Pt(s)|Br^{-}(0.010 M)|Br_{2}(l )||H^{+}(0.030 M)| H_{2}(g) (1 bar)|Pt(s)\).
- Q2.6: In the button cells widely used in watches and other devices the following reaction takes place: (Zn(s) + Ag_{2}O(s) + H_{2}O(l ) \rightarrow Zn^{2+}(aq) + 2Ag(s) + 2OH^{-}(aq)) Determine (\Delta_{r}G^{o}) and (E^{o}) for the reaction.
- Q2.6: In the button cells widely used in watches and other devices the following reaction takes place: \(Zn(s) + Ag_{2}O(s) + H_{2}O(l ) \rightarrow Zn^{2+}(aq) + 2Ag(s) + 2OH^{-}(aq)\) Determine \(\Delta_{r}G^{\circ}\) and \(E^{\circ}\) for the reaction.
- Q2.7: Define conductivity and molar conductivity for the solution of an electrolyte. Discuss their variation with concentration.
- Q2.7: Define conductivity and molar conductivity for the solution of an electrolyte. Discuss their variation with concentration.
- Q2.8: The conductivity of 0.20 M solution of KCl at 298 K is 0.0248 S (cm^{-1}). Calculate its molar conductivity.
- Q2.8: The conductivity of 0.20 M solution of KCl at 298 K is \(0.0248 S cm^{-1}\). Calculate its molar conductivity.
- Q2.9: The resistance of a conductivity cell containing 0.001M KCl solution at 298 K is 1500 (\Omega). What is the cell constant if conductivity of 0.001M KCl solution at 298 K is 0.146 (\times) (10^{-3}) S (cm^{-1}).
- Q2.9: The resistance of a conductivity cell containing 0.001M KCl solution at 298 K is \(1500 \Omega\). What is the cell constant if conductivity of 0.001M KCl solution at 298 K is \(0.146 \times 10^{-3} S cm^{-1}\).
CBSE Solutions for Class 12 Chemistry Electrochemistry
Chapters in CBSE - Class 12 Chemistry
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