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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 \(Fe^{2+} (aq)\).
Solution :
Step 1: Identification of Half-Cell Reactions
To determine the feasibility of the reaction between $Br_{2}(aq)$ and $Fe^{2+}(aq)$, we must identify the respective reduction half-reactions and their associated Standard Electrode Potentials ($E^{\circ}$) at $298\text{ K}$.
- Reduction half-reaction for Bromine: $Br_{2}(aq) + 2e^{-} \rightarrow 2Br^{-}(aq)$; $E^{\circ}_{Br_{2}/Br^{-}} = +1.09\text{ V}$
- Reduction half-reaction for Iron(II): $Fe^{3+}(aq) + e^{-} \rightarrow Fe^{2+}(aq)$; $E^{\circ}_{Fe^{3+}/Fe^{2+}} = +0.77\text{ V}$
[Per the electrochemical series, substances with higher positive reduction potentials act as stronger oxidizing agents.]
Step 2: Defining the Overall Cell Reaction
For the reaction to be feasible, the overall cell potential ($E^{\circ}_{cell}$) must be positive. This implies that the species with the higher reduction potential must undergo reduction, and the species with the lower reduction potential must undergo oxidation.
Given the standard potentials, $Br_{2}$ ($+1.09\text{ V}$) is a stronger oxidizing agent than $Fe^{3+}$ ($+0.77\text{ V}$). Therefore, $Br_{2}$ will be reduced, and $Fe^{2+}$ will be oxidized to $Fe^{3+}$.
The balanced cell reaction is:
$Br_{2}(aq) + 2Fe^{2+}(aq) \rightarrow 2Br^{-}(aq) + 2Fe^{3+}(aq)$
Step 3: Calculating Standard Cell Potential ($E^{\circ}_{cell}$)
The standard cell potential is calculated using the formula [By convention for electrochemical cells]:
$E^{\circ}_{cell} = E^{\circ}_{cathode} - E^{\circ}_{anode}$
Substituting the values:
$E^{\circ}_{cell} = E^{\circ}_{Br_{2}/Br^{-}} - E^{\circ}_{Fe^{3+}/Fe^{2+}}$
$E^{\circ}_{cell} = (+1.09\text{ V}) - (+0.77\text{ V})$
$E^{\circ}_{cell} = +0.32\text{ V}$
Step 4: Feasibility Analysis
According to the Gibbs Free Energy relation $\Delta G^{\circ} = -nFE^{\circ}_{cell}$, a reaction is spontaneous if $\Delta G^{\circ} < 0$. This condition is satisfied when $E^{\circ}_{cell} > 0$.
Since $E^{\circ}_{cell} = +0.32\text{ V} > 0$, the reaction is thermodynamically spontaneous under standard conditions.
Final Solution: The reaction between $Br_{2}(aq)$ and $Fe^{2+}(aq)$ is feasible, as the calculated standard cell potential $E^{\circ}_{cell} = +0.32\text{ V}$ is positive, indicating a spontaneous process.
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 \(Fe^{3+} (aq)\)
- Q2.18 (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 (iv): Predict the products of electrolysis in each of the following: (iv) An aqueous solution of (CuCl_{2}) 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
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