Electrochemistry carries 9 Marks in your CBSE Class 12 Board Exam. It is a "High Return on Investment" chapter. If you master these 50 questions, you cover 99% of the possible exam paper.
Bigyanbook Strategy: We have divided questions into "Levels" to help you climb the ladder of difficulty without stress. Click on a question to reveal the answer, analysis, and memory tricks!
📚 Syllabus: Full Unit 2
⏱️ Read Time: 25 Mins
🎯 Target: 70/70
🚀 Level 1: Rapid Fire (1 Mark VSA)
Clear your concepts. These questions are often asked in MCQs or 1-word answers.
1. Define Electrochemical Cell.
A device that converts chemical energy of a spontaneous redox reaction into electrical energy (Galvanic cell) or uses electrical energy to carry out a non-spontaneous reaction (Electrolytic cell).
2. What is the unit of Molar Conductivity ($\Lambda_m$)?
The unit is $S \ cm^2 \ mol^{-1}$ or $S \ m^2 \ mol^{-1}$.
💡 Tip: Don't forget the 'mol inverse'. Marks are cut for wrong units.
3. Define Cell Constant. What is its unit?
The ratio of the distance between electrodes ($l$) to the area of cross-section ($A$). Formula: $G^* = l/A$. Unit: $cm^{-1}$ or $m^{-1}$.
4. Why does the conductivity of a solution decrease with dilution?
Conductivity is the conductance of ions per unit volume. Upon dilution, the number of ions per unit volume decreases, hence conductivity decreases.
5. What is the Standard Electrode Potential of the Standard Hydrogen Electrode (SHE)?
It is arbitrarily taken as 0.00 V at all temperatures.
6. How does temperature affect the conductivity of metallic conductors vs electrolytic conductors?
- Metallic: Conductivity decreases as Temp increases (due to vibration of kernels).
- Electrolytic: Conductivity increases as Temp increases (due to increased mobility and dissociation).
7. Write the Nernst equation for a single electrode potential: $M^{n+} + ne^- \rightarrow M(s)$.
$E = E^{\circ} - \frac{RT}{nF} \ln \frac{1}{[M^{n+}]}$ or $E = E^{\circ} + \frac{0.059}{n} \log [M^{n+}]$
8. Can we store Copper Sulphate solution in a Zinc pot?
No. Zinc is more reactive ($E^{\circ} = -0.76V$) than Copper ($E^{\circ} = +0.34V$). Zn will displace Cu from the solution ($Zn + CuSO_4 \rightarrow ZnSO_4 + Cu$), creating holes in the pot.
💡 Trick: Lower Standard Potential = Higher Reactivity (Oxidation).
9. What is the relationship between $\Delta G^{\circ}$ and $E^{\circ}_{cell}$?
$\Delta G^{\circ} = -n F E^{\circ}_{cell}$.
10. What flows in the internal circuit of a Galvanic cell?
Ions flow in the internal circuit (via salt bridge), while electrons flow in the external circuit.
11. Define Limiting Molar Conductivity.
The molar conductivity of an electrolyte when the concentration approaches zero (Infinite Dilution). Denoted by $\Lambda^{\circ}_m$.
12. Which cell is used in hearing aids?
Mercury Cell (Button cell). It provides a constant voltage of 1.35V.
13. What is the efficiency of a Fuel Cell?
$\text{Efficiency} = \frac{\Delta G}{\Delta H} \times 100$. Theoretically 100%, practically ~70%.
14. Name the catalyst used in $H_2-O_2$ fuel cells.
Finely divided platinum or palladium incorporated into carbon electrodes.
15. What is the sign of $\Delta G$ for an electrolytic cell?
Positive. Electrolytic cells drive non-spontaneous reactions.
🧠 Level 2: Conceptual & Reasoning (2-3 Marks)
This is where most students get confused. Read the Bigyanbook Analysis carefully.
16. Why is Alternating Current (AC) used for measuring resistance of an electrolytic solution?
1. Direct Current (DC) causes electrolysis, which changes the concentration of the solution.
2. DC causes polarization at electrodes.
AC prevents electrolysis by constantly reversing polarity.
2. DC causes polarization at electrodes.
AC prevents electrolysis by constantly reversing polarity.
17. Explain the function of the Salt Bridge.
1. It completes the electrical circuit by allowing ions to flow.
2. It maintains electrical neutrality in the two half-cells by providing inert ions (like $K^+$, $NO_3^-$).
2. It maintains electrical neutrality in the two half-cells by providing inert ions (like $K^+$, $NO_3^-$).
18. State Faraday's Second Law of Electrolysis.
When the same quantity of electricity is passed through different electrolytes connected in series, the masses of substances produced at the electrodes are directly proportional to their equivalent masses (Equivalent Weight).
$\frac{W_1}{W_2} = \frac{E_1}{E_2}$
$\frac{W_1}{W_2} = \frac{E_1}{E_2}$
19. Suggest two materials other than hydrogen that can be used as fuels in fuel cells.
Methane ($CH_4$) and Methanol ($CH_3OH$).
20. Why does the voltage of a Mercury cell remain constant during its lifetime?
The overall cell reaction does not involve any ions in the solution whose concentration can change during its lifetime.
Reaction: $Zn(Hg) + HgO(s) \rightarrow ZnO(s) + Hg(l)$. All species are solids or liquids.
Reaction: $Zn(Hg) + HgO(s) \rightarrow ZnO(s) + Hg(l)$. All species are solids or liquids.
21. Predict the products of electrolysis of Aqueous NaCl.
At Cathode: $H_2$ gas (Reduction of water is preferred over Na+).
At Anode: $Cl_2$ gas (Oxidation of Cl- is preferred over water due to Overpotential).
In Solution: NaOH is formed.
At Anode: $Cl_2$ gas (Oxidation of Cl- is preferred over water due to Overpotential).
In Solution: NaOH is formed.
Analysis: This is a 5-star question. Remember that for Molten NaCl, product is Na metal. For Aqueous, it is H2 gas.
22. Define Electrochemical Equivalent ($Z$).
It is the mass of the substance deposited or liberated by passing 1 Coulomb of charge (or 1 Ampere current for 1 Second).
23. What is 'Overpotential' or 'Overvoltage'?
Some electrochemical processes (like the evolution of $O_2$ at the anode) are kinetically slow and require a potential higher than the theoretical value to occur. This extra potential is called Overpotential.
24. Differentiate between Metallic and Electrolytic conduction.
| Metallic | Electrolytic |
|---|---|
| Flow of electrons | Flow of ions |
| No matter transfer | Matter is transferred |
| Decreases with Temp | Increases with Temp |
25. Arrange Ag, Mg, K, and Cr in increasing order of reducing power given their standard potentials.
Given: $E^{\circ}$ values: $K (-2.93V), Mg (-2.37V), Cr (-0.74V), Ag (+0.80V)$.
Logic: More negative potential = stronger reducing agent.
Order: $Ag < Cr < Mg < K$.
Logic: More negative potential = stronger reducing agent.
Order: $Ag < Cr < Mg < K$.
🔥 Level 3: Numerical & Application (3-5 Marks)
Prepare your calculator (or log table). These questions carry the most weight.
26. Calculate $\Lambda^{\circ}_m$ for Acetic Acid using Kohlrausch Law.
Data given usually: $\lambda^{\circ}(H^+)$, $\lambda^{\circ}(Na^+)$, etc.
Formula: $\Lambda^{\circ}_{CH_3COOH} = \lambda^{\circ}_{H^+} + \lambda^{\circ}_{CH_3COO^-}$.
If salt data is given: $\Lambda^{\circ}_{CH_3COOH} = \Lambda^{\circ}_{CH_3COONa} + \Lambda^{\circ}_{HCl} - \Lambda^{\circ}_{NaCl}$.
Formula: $\Lambda^{\circ}_{CH_3COOH} = \lambda^{\circ}_{H^+} + \lambda^{\circ}_{CH_3COO^-}$.
If salt data is given: $\Lambda^{\circ}_{CH_3COOH} = \Lambda^{\circ}_{CH_3COONa} + \Lambda^{\circ}_{HCl} - \Lambda^{\circ}_{NaCl}$.
27. A solution of $CuSO_4$ is electrolyzed for 10 minutes with a current of 1.5 Amperes. What is the mass of copper deposited?
1. $Q = I \times t = 1.5 \times (10 \times 60) = 900 \ C$.
2. Reaction: $Cu^{2+} + 2e^- \rightarrow Cu$. ($n=2$).
3. 2F (2 $\times$ 96500 C) deposits 63.5g (Molar mass of Cu).
4. Mass = $\frac{63.5 \times 900}{2 \times 96500} = 0.296 \ g$.
2. Reaction: $Cu^{2+} + 2e^- \rightarrow Cu$. ($n=2$).
3. 2F (2 $\times$ 96500 C) deposits 63.5g (Molar mass of Cu).
4. Mass = $\frac{63.5 \times 900}{2 \times 96500} = 0.296 \ g$.
28. Calculate the emf of the cell: $Mg | Mg^{2+}(0.001M) || Cu^{2+}(0.0001M) | Cu$
Given $E^{\circ}_{cell} = 2.71V$.
Using Nernst Eq: $E = E^{\circ} - \frac{0.059}{2} \log \frac{[Mg^{2+}]}{[Cu^{2+}]}$
$E = 2.71 - 0.0295 \log \frac{10^{-3}}{10^{-4}}$
$E = 2.71 - 0.0295 \log 10 = 2.71 - 0.0295 = 2.68 V$.
Using Nernst Eq: $E = E^{\circ} - \frac{0.059}{2} \log \frac{[Mg^{2+}]}{[Cu^{2+}]}$
$E = 2.71 - 0.0295 \log \frac{10^{-3}}{10^{-4}}$
$E = 2.71 - 0.0295 \log 10 = 2.71 - 0.0295 = 2.68 V$.
💡 Analysis: Since Reactant concentration is lower than Product, EMF decreases.
29. Calculate $\Delta G^{\circ}$ and Equilibrium constant for $2Cr(s) + 3Cd^{2+} \rightarrow 2Cr^{3+} + 3Cd$.
$E^{\circ}_{cell} = +0.34V$. Electrons exchanged $n = 6$.
1. $\Delta G^{\circ} = -nFE^{\circ} = -6 \times 96500 \times 0.34 = -196860 \ J/mol$.
2. $\log K_c = \frac{nE^{\circ}}{0.059} = \frac{6 \times 0.34}{0.059} = 34.57$.
3. $K_c = \text{antilog}(34.57)$.
1. $\Delta G^{\circ} = -nFE^{\circ} = -6 \times 96500 \times 0.34 = -196860 \ J/mol$.
2. $\log K_c = \frac{nE^{\circ}}{0.059} = \frac{6 \times 0.34}{0.059} = 34.57$.
3. $K_c = \text{antilog}(34.57)$.
30. The resistance of a conductivity cell with 0.1M KCl is 100 $\Omega$. With 0.02M KCl, it is 520 $\Omega$. Calculate conductivity and molar conductivity of 0.02M solution.
(Conductivity of 0.1M KCl is 1.29 S/m).
Step 1: Find Cell Constant ($G^*$).
$G^* = \kappa \times R = 1.29 \times 100 = 129 \ m^{-1}$.
Step 2: Find $\kappa$ of 0.02M.
$\kappa = G^*/R = 129 / 520 = 0.248 \ S/m$.
Step 3: Find $\Lambda_m$.
$\Lambda_m = \frac{\kappa \times 1000}{Molarity}$ (Careful with units!).
Step 1: Find Cell Constant ($G^*$).
$G^* = \kappa \times R = 1.29 \times 100 = 129 \ m^{-1}$.
Step 2: Find $\kappa$ of 0.02M.
$\kappa = G^*/R = 129 / 520 = 0.248 \ S/m$.
Step 3: Find $\Lambda_m$.
$\Lambda_m = \frac{\kappa \times 1000}{Molarity}$ (Careful with units!).
31. How much charge is required for the reduction of 1 mol of $MnO_4^-$ to $Mn^{2+}$?
Change in Oxidation State: $+7 \rightarrow +2$.
Difference = 5 electrons.
Charge = $5F = 5 \times 96500 \ C = 482500 \ C$.
Difference = 5 electrons.
Charge = $5F = 5 \times 96500 \ C = 482500 \ C$.
32. How many hours does it take to reduce 3 mol of $Fe^{3+}$ to $Fe^{2+}$ with a 2.0 A current?
Reaction: $Fe^{3+} + e^- \rightarrow Fe^{2+}$. ($n=1$).
For 3 moles, we need 3 Faradays ($3 \times 96500$ C).
$Q = I \times t \Rightarrow t = Q/I = \frac{3 \times 96500}{2} = 144750 \ s$.
In hours: $144750 / 3600 = 40.2$ hours.
For 3 moles, we need 3 Faradays ($3 \times 96500$ C).
$Q = I \times t \Rightarrow t = Q/I = \frac{3 \times 96500}{2} = 144750 \ s$.
In hours: $144750 / 3600 = 40.2$ hours.
33. The conductivity of 0.001M acetic acid is $4.95 \times 10^{-5} S cm^{-1}$. Calculate dissociation constant ($K_a$) if $\Lambda^{\circ}_m = 390.5 S cm^2 mol^{-1}$.
1. Calc $\Lambda_m = \frac{1000 \times \kappa}{C} = \frac{1000 \times 4.95 \times 10^{-5}}{0.001} = 49.5$.
2. Calc degree of dissociation $\alpha = \Lambda_m / \Lambda^{\circ}_m = 49.5 / 390.5 \approx 0.127$.
3. $K_a = \frac{C \alpha^2}{1-\alpha}$. Substitute and solve.
2. Calc degree of dissociation $\alpha = \Lambda_m / \Lambda^{\circ}_m = 49.5 / 390.5 \approx 0.127$.
3. $K_a = \frac{C \alpha^2}{1-\alpha}$. Substitute and solve.
34. Predict spontaneity: $Fe^{3+}(aq) + I^-(aq) \rightarrow Fe^{2+}(aq) + I_2(s)$.
Given $E^{\circ}_{Fe^{3+}/Fe^{2+}} = 0.77V$, $E^{\circ}_{I_2/I^-} = 0.54V$.
$E^{\circ}_{cell} = E^{\circ}_{cat} - E^{\circ}_{an}$.
Reduction is Fe ($0.77$), Oxidation is I ($0.54$).
$E^{\circ}_{cell} = 0.77 - 0.54 = +0.23V$.
Since positive, reaction is spontaneous.
$E^{\circ}_{cell} = E^{\circ}_{cat} - E^{\circ}_{an}$.
Reduction is Fe ($0.77$), Oxidation is I ($0.54$).
$E^{\circ}_{cell} = 0.77 - 0.54 = +0.23V$.
Since positive, reaction is spontaneous.
35. Three electrolytic cells A, B, C containing $ZnSO_4$, $AgNO_3$, and $CuSO_4$ respectively are connected in series. A current of 1.5A was passed until 1.45g of Ag deposited. How long did current flow? What mass of Cu and Zn deposited?
Use Faraday's First Law for Time ($w=ZIt$).
Use Faraday's Second Law for Mass of Cu and Zn ($w_1/w_2 = E_1/E_2$).
Use Faraday's Second Law for Mass of Cu and Zn ($w_1/w_2 = E_1/E_2$).
🛡️ Level 4: Theory Vault (Batteries & Corrosion)
Rote memorization needed here. Use the mnemonics provided.
36. Write the anode and cathode reaction of a Dry Cell (Leclanche Cell).
Anode: $Zn(s) \rightarrow Zn^{2+} + 2e^-$.
Cathode: $MnO_2 + NH_4^+ + e^- \rightarrow MnO(OH) + NH_3$.
Cathode: $MnO_2 + NH_4^+ + e^- \rightarrow MnO(OH) + NH_3$.
37. Why does a dry cell become dead after some time even if not in use?
The acidic $NH_4Cl$ corrodes the zinc container.
38. Write the overall reaction of Lead Storage Battery during charging.
$2PbSO_4(s) + 2H_2O(l) \rightarrow Pb(s) + PbO_2(s) + 2H_2SO_4(aq)$.
(Just reverse the discharge reaction).
(Just reverse the discharge reaction).
39. Write the reactions occurring in a Ni-Cd cell.
Overall: $Cd(s) + 2Ni(OH)_3(s) \rightarrow CdO(s) + 2Ni(OH)_2(s) + H_2O(l)$.
40. What is Galvanization?
Coating iron with Zinc. Since Zinc ($E^{\circ}=-0.76V$) is more reactive than Iron ($E^{\circ}=-0.44V$), Zinc corrodes preferentially, protecting the iron (Sacrificial Protection).
41. How does $H^+$ ion concentration affect corrosion?
Higher $H^+$ concentration (Acidic medium) accelerates corrosion because $H^+$ ions are required for the reduction of Oxygen at the cathode of the corrosion cell.
42. What is cathodic protection?
Connecting the iron object to a more active metal (like Mg or Zn). The active metal becomes the anode and loses electrons, keeping iron as the cathode (protected).
43. Why are fuel cells preferred in space programs?
1. High efficiency/light weight.
2. The product is pure water, which astronauts can drink.
2. The product is pure water, which astronauts can drink.
44. Define SHE (Standard Hydrogen Electrode) construction.
It consists of a Platinum wire coated with Platinum black, dipped in 1M $H^+$ solution, with pure $H_2$ gas bubbled at 1 bar pressure.
45. What is the role of Platinum in SHE?
It provides a surface for the oxidation or reduction of Hydrogen and conducts electrons (Inert Electrode).
46. Why is fluorine the strongest oxidizing agent?
It has the highest positive Standard Electrode Potential ($E^{\circ} = +2.87V$), meaning it has the highest tendency to get reduced (gain electrons).
47. Can $E^{\circ}_{cell}$ be negative?
No. By definition, a Galvanic cell utilizes a spontaneous reaction. If calculated potential is negative, the reaction is non-spontaneous and the cell won't work (or it's an electrolytic cell).
48. Why is conductivity of water not zero?
Even pure water self-ionizes slightly into $H^+$ and $OH^-$. Also, dissolved $CO_2$ from air forms Carbonic acid, contributing ions.
49. Variation of $\Lambda_m$ for Strong vs Weak electrolytes (Graphical).
Strong (KCl): Linear small increase with dilution (Debye-Huckel Onsager eq).
Weak (Acetic Acid): Steep curve increase at infinite dilution due to increase in degree of dissociation.
Weak (Acetic Acid): Steep curve increase at infinite dilution due to increase in degree of dissociation.
50. Calculate the potential of Hydrogen electrode at pH = 10.
Reaction: $H^+ + e^- \rightarrow \frac{1}{2}H_2$.
pH = 10 means $[H^+] = 10^{-10} M$.
$E = E^{\circ} - 0.059 \log \frac{1}{[H^+]} = 0 - 0.059 (\text{pH})$.
$E = -0.059 \times 10 = -0.59 V$.
pH = 10 means $[H^+] = 10^{-10} M$.
$E = E^{\circ} - 0.059 \log \frac{1}{[H^+]} = 0 - 0.059 (\text{pH})$.
$E = -0.059 \times 10 = -0.59 V$.
💡 Super Trick: For Hydrogen electrode, $E = -0.059 \times pH$. Memorize this!
Bigyanbook Final Advice
You have just reviewed the entire chapter. For the exam:
- Practice the Unit conversions in Conductivity numericals (cm to m).
- Draw the Batteries diagram at least once.
- Don't panic on the calculation. Use approximations (e.g., 96500 is approx $10^5$ for rough check).
All the best for your Boards!