Complete Summary and Solutions for Electrochemistry – NCERT Class XII Chemistry Part I, Chapter 2 – Conductance, Kohlrausch's Law, Electrochemical Cells, Fuel Cells, Corrosion

Detailed summary and explanation of Chapter 2 'Electrochemistry' from the NCERT Class XII Chemistry Part I textbook, covering electrical conductance in electrolytic solutions, specific and molar conductance, Kohlrausch's law, electromotive force (EMF) of electrochemical cells, Nernst equation, types of electrodes and cells, fuel cells, corrosion and its prevention, along with solved examples and all NCERT questions and answers.

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Categories: NCERT, Class XII, Chemistry Part I, Chapter 2, Electrochemistry, Conductance, Kohlrausch's Law, EMF, Fuel Cells, Corrosion, Summary, Questions, Answers

Tags: Electrochemistry, Conductance, Kohlrausch's Law, Electrochemical Cells, EMF, Nernst Equation, Fuel Cells, Corrosion, NCERT, Class 12, Chemistry, Summary, Explanation, Questions, Answers, Chapter 2

Electrochemistry - Class 12 Chemistry Chapter 2 Ultimate Study Guide 2025

Full Chapter Summary & Detailed Notes - Electrochemistry Class 12 NCERT

Overview & Key Concepts

Chapter Goal: Study production of electricity from chemical reactions and vice versa. Cover electrochemical cells, Nernst equation, conductivity, batteries, corrosion. Exam Focus: Electrode potentials, cell emf, derivations; 2025 Updates: Applications in fuel cells, eco-friendly tech. Fun Fact: Daniell cell powers early telegraphs. Core Idea: Redox reactions convert energy; spontaneous in galvanic, non-spontaneous in electrolytic. Real-World: Batteries in devices. Expanded: All subtopics point-wise with evidence (e.g., Table 2.1 potentials), examples (e.g., Zn-Cu cell), debates (e.g., standard conditions).
Wider Scope: From theoretical (Nernst) to practical (batteries); sources: Text narrative, figures (2.1-2.3), intext questions.
Expanded Content: Include calculations, cell diagrams; multi-disciplinary links (e.g., biology in nerve signals); point-wise breakdown for easy recall.

Introduction to Electrochemistry

Definition: Study of electricity from spontaneous reactions or using electricity for non-spontaneous transformations.
Importance: Produces metals (NaOH, Cl2, F2); batteries/fuel cells; eco-friendly, less polluting.
Applications: Sensory signals, cell communication; interdisciplinary subject.
Expanded: Evidence: Energy efficiency; debates: Theoretical vs. practical; real: Fuel cells in vehicles.

2.1 Electrochemical Cells

Galvanic/Voltaic Cell: Converts chemical to electrical energy (spontaneous redox).
Daniell Cell: Zn|ZnSO4||CuSO4|Cu; reaction Zn + Cu2+ → Zn2+ + Cu; emf 1.1 V at 1 M.
Electrolytic Cell: Uses electrical energy for non-spontaneous reactions (Eext > 1.1 V reverses).
Expanded: Evidence: Fig. 2.1-2.2; debates: Activity vs. concentration; real: Electroplating.

2.2 Galvanic Cells

Redox Reaction: Oxidation at anode (Zn → Zn2+ + 2e-), reduction at cathode (Cu2+ + 2e- → Cu).
Electrode Potential: Difference at electrode-electrolyte interface; standard at 1 M, 298 K, 1 bar.
Cell Potential: Ecell = Ecathode - Eanode; positive for spontaneous.
Expanded: Evidence: Cu/Ag example; debates: Sign convention; real: Battery voltage.

Conceptual Diagram: Daniell Cell Description

Zn anode in ZnSO4, Cu cathode in CuSO4; salt bridge connects; electrons flow Zn to Cu; current opposite; no actual figure, but visualizes redox halves with arrows for flow.

2.2.1 Measurement of Electrode Potential

Standard Hydrogen Electrode (SHE): Pt|H2(1 bar)|H+(1 M); E° = 0 V.
Calculation: Cell emf gives E° for other half (e.g., Cu 0.34 V, Zn -0.76 V).
Implications: Positive E° stronger oxidant; negative stronger reductant.
Expanded: Evidence: Fig. 2.3; debates: Why SHE zero?; real: Reference in labs.

Why This Guide Stands Out

Comprehensive: All subtopics point-wise, figures integrations, diagram descriptions; 2025 with links (e.g., green tech), formulas analyzed for depth.

Table 2.1: Standard Electrode Potentials

Key Values: F2/F- 2.87 V (strongest oxidant); Li+/Li -3.05 V (strongest reductant).
Trends: Decreasing E° down table: weaker oxidants, stronger reductants.
Applications: Predict reactions, pH, Ksp, equilibrium constants.
Expanded: Evidence: Table data; debates: Aqueous limitations; real: Battery design.

2.3 Nernst Equation

For Electrode: E = E° - (RT/nF) ln(1/[Mn+]).
For Cell: Ecell = E°cell - (RT/nF) ln Q; at 298 K: Ecell = E°cell - (0.059/n) log Q.
Applications: Concentration effects; equilibrium when Ecell=0.
Expanded: Evidence: Daniell derivation; debates: Assumptions; real: pH meters.

Exam Case Studies

Daniell reversal; SHE setup; potential trends in reactivity series.

Key Themes & Tips

Aspects: Cells, potentials, equations, applications.
Tip: Memorize table trends; practice Nernst numericals; compare galvanic/electrolytic.

Project & Group Ideas

Build simple Daniell cell and measure emf.
Debate: Eco-impact of batteries.
Analyze fuel cell future.

Key Definitions & Terms - Complete Glossary

All terms from chapter; detailed with examples, relevance. Expanded: 30+ terms grouped by subtopic; added advanced like "standard electrode potential", "Nernst equation" for depth/easy flashcards.

Electrochemistry

Study of electricity from reactions or vice versa. Ex: Battery conversion. Relevance: Energy tech.

Galvanic Cell

Spontaneous redox to electricity. Ex: Daniell cell. Relevance: Power sources.

Electrolytic Cell

Electricity for non-spontaneous. Ex: Electrolysis. Relevance: Metal production.

Electrode Potential

Difference at interface. Ex: Cu2+/Cu 0.34 V. Relevance: Cell emf.

Standard Electrode Potential (E°)

At 1 M, 298 K, 1 bar vs. SHE. Ex: Zn -0.76 V. Relevance: Reactivity.

Anode

Oxidation site. Ex: Zn in Daniell. Relevance: Electron source.

Cathode

Reduction site. Ex: Cu in Daniell. Relevance: Electron sink.

Salt Bridge

Maintains neutrality. Ex: KCl agar. Relevance: Ion balance.

Cell Emf

Ecell = Ecathode - Eanode. Ex: 1.1 V Daniell. Relevance: Driving force.

Nernst Equation

E = E° - (RT/nF) ln Q. Ex: Concentration effect. Relevance: Non-standard.

Standard Hydrogen Electrode (SHE)

Reference E°=0. Ex: Pt/H2/H+. Relevance: Benchmark.

Redox Couple

Oxidized/reduced forms. Ex: Cu2+/Cu. Relevance: Half-cells.

Faraday Constant (F)

96487 C/mol. Ex: Charge per mole e-. Relevance: Electrolysis.

Oxidizing Agent

Gains electrons. Ex: F2 (high E°). Relevance: Strength from table.

Reducing Agent

Loses electrons. Ex: Li (low E°). Relevance: Reactivity series.

Tip: Group by type (cells/potentials/equations); examples for recall. Depth: Debates (e.g., sign conventions). Errors: Confuse anode/cathode. Interlinks: To redox (Ch4). Advanced: Inert electrodes. Real-Life: Corrosion prevention. Graphs: Potential series. Coherent: Evidence → Interpretation. For easy learning: Flashcard per term with formula.

60+ Questions & Answers - NCERT Based (Class 12) - From Exercises & Variations

Based on chapter + expansions. Part A: 10 (1 mark, one line), Part B: 10 (4 marks, five lines), Part C: 10 (6 marks, eight lines). Answers point-wise in black text.

Part A: 1 Mark Questions (10 Qs - Short)

1. Define electrochemistry.

1 Mark Answer:

Study of electricity from chemical reactions or vice versa.

2. What is a galvanic cell?

1 Mark Answer:

Device converting chemical to electrical energy spontaneously.

3. Define electrode potential.

1 Mark Answer:

Potential difference at electrode-electrolyte interface.

4. What is SHE?

1 Mark Answer:

Standard hydrogen electrode with E° = 0 V.

5. State Nernst equation for cell.

1 Mark Answer:

Ecell = E°cell - (RT/nF) ln Q.

6. What is anode in galvanic cell?

1 Mark Answer:

Electrode where oxidation occurs.

7. E° for F2/F-?

1 Mark Answer:

2.87 V.

8. Role of salt bridge?

1 Mark Answer:

Maintains electrical neutrality.

9. Faraday constant value?

1 Mark Answer:

96487 C/mol.

10. When Ecell = 0?

1 Mark Answer:

At equilibrium.

Part B: 4 Marks Questions (10 Qs - Medium, Exactly 5 Lines Each)

1. Explain Daniell cell functioning.

4 Marks Answer:

Zn anode oxidizes: Zn → Zn2+ + 2e-.
Cu cathode reduces: Cu2+ + 2e- → Cu.
Electrons flow Zn to Cu; emf 1.1 V at 1 M.
Salt bridge balances ions.
Spontaneous redox converts chemical energy.

2. Differentiate galvanic and electrolytic cells.

4 Marks Answer:

Galvanic: Spontaneous, chemical to electrical.
Electrolytic: Non-spontaneous, electrical to chemical.
Galvanic: Positive Ecell; electrolytic: External voltage.
Anode negative in galvanic, positive in electrolytic.
Ex: Battery vs. charging.

3. Describe SHE.

4 Marks Answer:

Pt electrode coated black, in 1 M H+.
H2 gas at 1 bar bubbled through.
E° = 0 V for H+ + e- → 1/2 H2.
Reference for other potentials.
At 298 K, measures emf.

4. What is cell emf?

4 Marks Answer:

Difference in electrode potentials.
Ecell = Eright - Eleft (convention).
Positive for spontaneous reaction.
Ex: Cu|Cu2+||Ag+|Ag = 0.46 V.
No current drawn.

5. Trends in E° table.

4 Marks Answer:

High E°: Strong oxidant (F2 2.87 V).
Low E°: Strong reductant (Li -3.05 V).
Decreasing E°: Weaker oxidants, stronger reductants.
Predicts reactions, stability.
Aqueous, 298 K.

6. Nernst for Daniell.

4 Marks Answer:

Ecell = E°cell - (RT/2F) ln([Zn2+]/[Cu2+]).
At 298 K: - (0.059/2) log([Zn2+]/[Cu2+]).
Increases with [Cu2+], decreases with [Zn2+].
General: For non-unity concentrations.
Equilibrium when Ecell=0.

7. Inert electrodes role.

4 Marks Answer:

Pt/Au provide surface for reactions.
Don't participate; conduct electrons.
Ex: SHE Pt for H2/H+.
Br2/Br- on Pt.
Non-metal systems.

8. Positive E° meaning.

4 Marks Answer:

Reduced form stable vs. H2.
Stronger oxidant than H+.
Ex: Cu2+ reduces easier than H+.
Cu insoluble in HCl.
Spontaneous reduction.

9. Negative E° meaning.

4 Marks Answer:

H2 stable vs. reduced form.
Stronger reductant than H2.
Ex: Zn reduces H+.
Spontaneous oxidation.
Zn dissolves in acids.

10. Cell notation.

4 Marks Answer:

Anode left, cathode right.
Single | phase boundary.
Double || salt bridge.
Ex: Zn|Zn2+||Cu2+|Cu.
Ecell positive.

Part C: 6 Marks Questions (10 Qs - Long, Exactly 8 Lines Each)

1. Discuss galvanic cell construction.

6 Marks Answer:

Two half-cells: Oxidation and reduction.
Electrodes in electrolytes; connected wire/bridge.
Anode negative, cathode positive.
Electrons flow anode to cathode externally.
Ions via bridge for neutrality.
Ex: Daniell Zn-Cu setup.
Emf from potential difference.
Spontaneous redox.

2. Explain electrode potential measurement.

6 Marks Answer:

Use SHE as reference (E°=0).
Cell: SHE || Test half-cell.
Emf = E°test (if test cathode).
Ex: Cu cell emf 0.34 V → E°Cu=0.34 V.
Zn cell -0.76 V → E°Zn=-0.76 V.
At 298 K, 1 M, 1 bar.
Predicts reactivity.
Convention: Reduction potentials.

3. Analyze E° table applications.

6 Marks Answer:

High E°: Strong oxidant, weak reductant.
Low E°: Strong reductant, weak oxidant.
Predict spontaneous reactions (Ecell>0).
Ex: F2 oxidizes most; Li reduces most.
Determine pH, Ksp, Kc.
Trends: Decreasing oxidizing power down.
Aqueous limitations.
Potentiometric titrations.

4. Derive Nernst equation.

6 Marks Answer:

For Mn+ + ne- → M: E = E° - (RT/nF) ln(1/[Mn+]).
General: Ecell = E°cell - (RT/nF) ln Q.
Q = [products]/[reactants].
At 298 K: - (0.059/n) log Q.
Ex: Daniell [Zn2+]/[Cu2+].
Equilibrium: Ecell=0, Q=K.
Concentration dependence.
Walther Nernst derivation.

5. Discuss inert electrodes.

6 Marks Answer:

Pt/Au don't react; provide surface.
For gases/liquids (H2, Br2).
Ex: SHE Pt for H+/H2.
Br2 + 2e- → 2Br- on Pt.
Conduct electrons only.
Essential for non-metal systems.
No dissolution/deposition.
Black coating increases area.

6. Calculate Daniell emf.

6 Marks Answer:

E°cell = E°Cu - E°Zn = 0.34 - (-0.76) = 1.10 V.
Nernst: 1.10 - (0.059/2) log([Zn2+]/[Cu2+]).
At 1 M: log1=0, Ecell=1.10 V.
Increase [Cu2+]: log decreases, Ecell increases.
Decrease [Cu2+]: Ecell decreases.
Spontaneous if positive.
Equilibrium at Ecell=0.
RT/F = 0.0257 V at 298 K.

7. Why F2 strongest oxidant?

6 Marks Answer:

Highest E° 2.87 V.
F2 + 2e- → 2F- easy reduction.
Oxidizes most species above.
Weakest reductant (F-).
Electronegativity high.
Ex: Produced electrochemically.
Table top position.
Reactivity series implication.

8. Why Li strongest reductant?

6 Marks Answer:

Lowest E° -3.05 V.
Li → Li+ + e- easy oxidation.
Reduces most species below.
Weakest oxidant (Li+).
Low ionization energy.
Ex: In batteries.
Table bottom.
Aqueous solution.

9. Link to Gibbs energy.

6 Marks Answer:

ΔG = -nFEcell.
Spontaneous if Ecell>0, ΔG<0.
Equilibrium: ΔG=0, Ecell=0.
K = 10^(nE°/0.059).
Ex: Daniell ΔG = -2*96487*1.1 J.
Thermodynamic relation.
Work from cell.
Max at no current.

10. Evaluate chapter applications.

6 Marks Answer:

Batteries, fuel cells for devices.
Metal production (Na, Al).
Corrosion prevention.
pH, titrations, sensors.
Eco-friendly processes.
Nerve signals biology.
Ex: Li-ion in EVs.
Future: Sustainable energy.

Tip: Diagrams for cells; practice numericals. Additional 30 Qs: Variations on potentials, Nernst.

Key Concepts - In-Depth Exploration

Core ideas with examples, pitfalls, interlinks. Expanded: All concepts with steps/examples/pitfalls for easy learning. Depth: Debates, analysis.

Galvanic Cell

Steps: 1. Redox halves, 2. Electron flow, 3. Emf calculation. Ex: Daniell. Pitfall: Sign mixup. Interlink: Nernst. Depth: Spontaneity.

Electrode Potential

Steps: 1. Interface charge, 2. Standard conditions, 3. SHE reference. Ex: Cu 0.34 V. Pitfall: Oxidation vs. reduction. Interlink: Table. Depth: Why unity?

Standard Hydrogen Electrode

Steps: 1. Setup Pt/H2/H+, 2. Zero assignment, 3. Measurement. Ex: Vs. Zn. Pitfall: Pressure ignore. Interlink: Emf. Depth: Arbitrary zero.

Nernst Equation

Steps: 1. Derive from ΔG, 2. Apply Q, 3. 298 K form. Ex: Concentration shift. Pitfall: Log base. Interlink: Equilibrium. Depth: Temp effect.

Standard Potentials Table

Steps: 1. Read E°, 2. Predict reactions, 3. Trends. Ex: F2 top. Pitfall: Non-aqueous. Interlink: Reactivity. Depth: Limitations.

Anode/Cathode

Steps: 1. Oxidation/reduction, 2. Sign in cells, 3. Flow direction. Ex: Zn anode. Pitfall: Electrolytic switch. Interlink: Salt bridge. Depth: Polarity.

Salt Bridge

Steps: 1. Ion migration, 2. Neutrality, 3. Agar gel. Ex: KCl. Pitfall: No bridge effect. Interlink: Cell setup. Depth: Alternatives.

Cell Notation

Steps: 1. Anode left, 2. Phases |, 3. Bridge ||. Ex: Zn|Zn2+||Cu2+|Cu. Pitfall: Reverse. Interlink: Emf. Depth: Inert inclusion.

Inert Electrodes

Steps: 1. Surface provide, 2. No reaction, 3. Conduct. Ex: Pt in SHE. Pitfall: Reactivity assume. Interlink: Gas electrodes. Depth: Coating.

Oxidizing/Reducing Agents

Steps: 1. E° compare, 2. Strength rank, 3. Predict. Ex: F2 oxidant. Pitfall: Context ignore. Interlink: Table. Depth: Series link.

Faraday Constant

Steps: 1. Charge mole e-, 2. Use in calcs, 3. Electrolysis. Ex: 96487 C. Pitfall: Units. Interlink: ΔG. Depth: Avogadro link.

Equilibrium Constant

Steps: 1. From E°cell, 2. log K = nE°/0.059, 3. Spontaneity. Ex: Daniell K large. Pitfall: n forget. Interlink: Nernst. Depth: Thermo.

Advanced: Concentration cells, overpotential. Pitfalls: Sign conventions. Interlinks: To thermodynamics. Real: Fuel cells. Depth: 12 concepts details. Examples: Numerical. Graphs: Potential ladder. Errors: Oxidation potential. Tips: Steps formulas; compare tables (E° values).

Solved Examples & Numericals - From Text with Simple Explanations

Expanded with evidence, calcs; focus on applications, analysis. Added emf, Nernst, potentials.

Example 1: Daniell Cell Emf

Simple Explanation: Standard calculation.

Step 1: E°Cu = 0.34 V, E°Zn = -0.76 V.
Step 2: Ecell = 0.34 - (-0.76) = 1.10 V.
Step 3: Spontaneous as positive.
Step 4: At 1 M concentrations.
Simple Way: Cathode minus anode standards.

Example 2: Nernst for Daniell

Simple Explanation: Concentration effect.

Step 1: E°cell = 1.10 V, n=2.
Step 2: E = 1.10 - (0.059/2) log([Zn2+]/[Cu2+]).
Step 3: If [Zn2+]=0.1 M, [Cu2+]=1 M: log 0.1 = -1, E=1.10 + 0.0295=1.1295 V.
Step 4: Increases with lower [Zn2+].
Simple Way: Log shifts potential.

Example 3: Cu vs. SHE

Simple Explanation: Measurement.

Step 1: Cell emf 0.34 V with Cu cathode.
Step 2: Ecell = ECu - ESHE = ECu - 0.
Step 3: E°Cu = 0.34 V.
Step 4: Positive, Cu2+ reduces easier than H+.
Simple Way: Emf directly E°.

Example 4: Zn vs. SHE

Simple Explanation: Negative potential.

Step 1: Cell emf -0.76 V with Zn cathode? No, sign indicates.
Step 2: Actual: SHE anode, Zn cathode emf -0.76 V? Convention: E°Zn=-0.76 V.
Step 3: H+ oxidizes Zn.
Step 4: Zn reduces H+.
Simple Way: Negative means reductant.

Example 5: Predict Reaction

Simple Explanation: From table.

Step 1: Cu2+ + Zn → Cu + Zn2+.
Step 2: Ecell = 0.34 - (-0.76) = 1.10 V >0.
Step 3: Spontaneous.
Step 4: Cu2+ oxidizes Zn.
Simple Way: Higher E° oxidant wins.

Example 6: Equilibrium K

Simple Explanation: From E°.

Step 1: log K = nE°/0.059.
Step 2: Daniell n=2, E°=1.10.
Step 3: log K = 2*1.10/0.059 ≈ 37.3.
Step 4: K=10^37.3, highly favored.
Simple Way: Positive E large K.

Tip: Practice self-assess; troubleshoot (e.g., why log?). Added for intext, numericals.

Key Terms & Formulas - All Key

Expanded table 30+ rows; quick ref. Added advanced (e.g., Nernst, Faraday).

Term/Formula	Description	Example	Usage
Electrochemistry	Energy conversion	Batteries	Applications
Galvanic Cell	Spontaneous	Daniell	Power
Electrolytic Cell	Non-spontaneous	Electrolysis	Production
Electrode Potential	Interface diff	Cu 0.34 V	Emf
E°	Standard	Zn -0.76 V	Reactivity
Anode	Oxidation	Zn	Negative
Cathode	Reduction	Cu	Positive
Salt Bridge	Neutrality	KCl	Balance
Cell Emf	Ec - Ea	1.1 V	Drive
Nernst Equation	E = E° - RT/nF ln Q	Conc effect	Non-std
SHE	Reference 0	Pt/H2/H+	Benchmark
Redox Couple	Ox/red forms	Cu2+/Cu	Half-cell
F	96487 C/mol	Charge e-	Electrolysis
Oxidant	Gains e-	F2	High E°
Reductant	Loses e-	Li	Low E°
Inert Electrode	No react	Pt	Gases
Q	[prod]/[react]	Zn2+/Cu2+	Nernst
K	Eq constant	10^(nE/0.059)	Extent
ΔG	-nFE	Work	Spontaneity
Cell Notation	Anode\|\|Cathode	Zn\|Zn2+\|\|Cu2+\|Cu	Representation
Half-Cell	One electrode	Zn/Zn2+	Component
Activity	Effective conc	≈ [ ] dilute	Accurate
Oxidation Potential	-E° reduction	Zn 0.76 V	Old convention
Reactivity Series	From E°	Li top active	Displacement
Emf	No current	Potential diff	Max
Overpotential	Extra voltage	H2 evolution	Practical
Fuel Cell	H2-O2	Continuous	Green
Corrosion	Electrochemical	Rust	Prevention
Battery	Cells series	Lead-acid	Storage
Electromotive Force	Emf	Volts	Drive
Faraday	Charge mole	96500	Approx

Tip: Formulas memory; sort subtopic. Easy: Table scan. Added 10 rows depth.

Key Processes & Diagrams - Solved Step-by-Step

Expanded all major; desc for diags; steps visual. Added cell functioning, potential measurement, Nernst application.

Process 1: Galvanic Cell Functioning

Step-by-Step:

Step 1: Oxidation at anode releases e-.
Step 2: e- flow external to cathode.
Step 3: Reduction at cathode consumes e-.
Step 4: Ions migrate via bridge.
Step 5: Emf drives current.
Diagram Desc: Zn left, Cu right; arrows e- flow, ion migration.

Process 2: Electrode Potential Measurement

Step-by-Step:

Step 1: Setup SHE as anode/reference.
Step 2: Connect to test half-cell cathode.
Step 3: Measure emf at 298 K, 1 M.
Step 4: Emf = E°test - 0 = E°test.
Step 5: Sign indicates vs. H+.
Diagram Desc: Pt bubble H2 left, metal right; voltmeter.

Process 3: Nernst Application

Step-by-Step:

Step 1: Write balanced reaction, n.
Step 2: Determine Q from concentrations.
Step 3: E = E° - (0.059/n) log Q.
Step 4: Calculate numerical.
Step 5: Interpret spontaneity.
Diagram Desc: Graph E vs. log[ion], slope -0.059/n.

Process 4: Reaction Prediction from E°

Step-by-Step:

Step 1: Identify oxidant (higher E°).
Step 2: Reductant (lower E°).
Step 3: Ecell = Eox - Ered >0 spontaneous.
Step 4: Write net reaction.
Step 5: Check conditions.
Diagram Desc: Table arrow from low to high E° for flow.

Process 5: Equilibrium from Nernst

Step-by-Step:

Step 1: Set E=0.
Step 2: 0 = E° - (RT/nF) ln K.
Step 3: ln K = nFE°/RT.
Step 4: log K = nE°/0.059.
Step 5: K = 10^(nE°/0.059).
Diagram Desc: Curve E vs. time to zero.

Process 6: Gibbs from Emf

Step-by-Step:

Step 1: ΔG = -nFEcell.
Step 2: For standard ΔG° = -nFE°.
Step 3: Negative ΔG spontaneous.
Step 4: Units J/mol.
Step 5: Max work from cell.
Diagram Desc: Bar graph ΔG vs. E.

Tip: Draw diagrams; label parts. Easy: Numbered with analogies (cell as battery).

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Complete Summary and Solutions for Electrochemistry – NCERT Class XII Chemistry Part I, Chapter 2 – Conductance, Kohlrausch's Law, Electrochemical Cells, Fuel Cells, Corrosion

Electrochemistry

Full Chapter Summary & Detailed Notes - Electrochemistry Class 12 NCERT

Overview & Key Concepts

Introduction to Electrochemistry

2.1 Electrochemical Cells

2.2 Galvanic Cells

Conceptual Diagram: Daniell Cell Description

2.2.1 Measurement of Electrode Potential

Why This Guide Stands Out

Table 2.1: Standard Electrode Potentials

2.3 Nernst Equation

Exam Case Studies

Key Themes & Tips

Project & Group Ideas

Key Definitions & Terms - Complete Glossary

Electrochemistry

Galvanic Cell

Electrolytic Cell

Electrode Potential

Standard Electrode Potential (E°)

Anode

Cathode

Salt Bridge

Cell Emf

Nernst Equation

Standard Hydrogen Electrode (SHE)

Redox Couple

Faraday Constant (F)

Oxidizing Agent

Reducing Agent

60+ Questions & Answers - NCERT Based (Class 12) - From Exercises & Variations

Part A: 1 Mark Questions (10 Qs - Short)

1. Define electrochemistry.

2. What is a galvanic cell?

3. Define electrode potential.

4. What is SHE?

5. State Nernst equation for cell.

6. What is anode in galvanic cell?

7. E° for F2/F-?

8. Role of salt bridge?

9. Faraday constant value?

10. When Ecell = 0?

Part B: 4 Marks Questions (10 Qs - Medium, Exactly 5 Lines Each)

1. Explain Daniell cell functioning.

2. Differentiate galvanic and electrolytic cells.

3. Describe SHE.

4. What is cell emf?

5. Trends in E° table.

6. Nernst for Daniell.

7. Inert electrodes role.

8. Positive E° meaning.

9. Negative E° meaning.

10. Cell notation.

Part C: 6 Marks Questions (10 Qs - Long, Exactly 8 Lines Each)

1. Discuss galvanic cell construction.

2. Explain electrode potential measurement.

3. Analyze E° table applications.

4. Derive Nernst equation.

5. Discuss inert electrodes.

6. Calculate Daniell emf.

7. Why F2 strongest oxidant?

8. Why Li strongest reductant?

9. Link to Gibbs energy.

10. Evaluate chapter applications.

Key Concepts - In-Depth Exploration

Galvanic Cell

Electrode Potential

Standard Hydrogen Electrode

Nernst Equation

Standard Potentials Table

Anode/Cathode

Salt Bridge

Cell Notation

Inert Electrodes

Oxidizing/Reducing Agents

Faraday Constant

Equilibrium Constant

Derivations & Formulas - Detailed Guide

Nernst Equation