CHAPTER 2 – ELECTROCHEMISTRY IMPORTANT QUESTION BANK

MCQs + Fill in the Blanks + Assertion Reason + Very Short Questions

Assertion Reason + Very Short Questions

SECTION A: MULTIPLE CHOICE QUESTIONS (MCQs)

1. In a galvanic cell, conversion takes place from:

A. Electrical energy to chemical energy
B. Chemical energy to electrical energy
C. Heat energy to electrical energy
D. Light energy to chemical energy

Answer: B

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2. Oxidation occurs at:

A. Cathode
B. Salt bridge
C. Anode
D. Electrolyte

Answer: C

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3. In a galvanic cell, electrons flow from:

A. Cathode to anode
B. Anode to cathode
C. Salt bridge to cathode
D. Electrolyte to electrode

Answer: B

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4. The standard electrode potential of Standard Hydrogen Electrode is:

A. 1 V
B. 0 V
C. –1 V
D. 0.5 V

Answer: B

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5. The unit of conductivity is:

A. Ω
B. S m⁻¹
C. S⁻¹
D. mol m⁻¹

Answer: B

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6. The function of a salt bridge is:

A. To increase resistance
B. To provide electrons
C. To maintain electrical neutrality
D. To stop ion movement

Answer: C

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7. Cell potential is calculated by:

A. Anode − Cathode
B. Cathode − Anode
C. Cathode + Anode
D. Anode × Cathode

Answer: B

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8. A positive value of E°cell indicates:

A. Non-spontaneous reaction
B. Equilibrium condition
C. Spontaneous reaction
D. No reaction

Answer: C

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9. Nernst equation is used to calculate:

A. Atomic mass
B. Electrode potential under non-standard conditions
C. Molecular mass
D. Density

Answer: B

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10. The value of Faraday constant is approximately:

A. 965 C mol⁻¹
B. 9650 C mol⁻¹
C. 96500 C mol⁻¹
D. 9.65 C mol⁻¹

Answer: C

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11. Conductance is the reciprocal of:

A. Conductivity
B. Resistance
C. Resistivity
D. Cell constant

Answer: B

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12. On dilution, conductivity of electrolyte solution:

A. Increases
B. Decreases
C. Remains constant
D. Becomes zero

Answer: B

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13. On dilution, molar conductivity:

A. Decreases
B. Increases
C. Remains unchanged
D. Becomes zero

Answer: B

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14. Strong electrolytes are:

A. Completely ionised
B. Not ionised
C. Weakly ionised
D. Insoluble

Answer: A

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15. Which law explains independent migration of ions?

A. Faraday’s law
B. Ohm’s law
C. Kohlrausch’s law
D. Henry’s law

Answer: C

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16. Weak electrolytes show a large increase in molar conductivity on dilution because:

A. Temperature increases
B. Ionisation increases
C. Resistance increases
D. Volume decreases

Answer: B

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17. In electrolysis, reduction occurs at:

A. Anode
B. Cathode
C. Salt bridge
D. Electrolyte

Answer: B

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18. Faraday’s first law relates deposited mass with:

A. Resistance
B. Voltage
C. Quantity of electricity
D. Temperature

Answer: C

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19. The formula for mass deposited during electrolysis is:

A. m = IR
B. m = MIt/nF
C. m = nF/MI
D. m = VIt

Answer: B

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20. Rust is chemically:

A. FeO
B. Fe₂O₃
C. Fe₂O₃·xH₂O
D. Fe(OH)₂

Answer: C

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21. A rechargeable battery is:

A. Dry cell
B. Mercury cell
C. Lead storage battery
D. Primary cell

Answer: C

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22. Hydrogen-oxygen fuel cell produces:

A. CO₂
B. H₂O
C. O₂
D. H₂O₂

Answer: B

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23. In a lead storage battery, electrolyte used is:

A. HCl
B. NaOH
C. H₂SO₄
D. KOH

Answer: C

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24. AC current is used in conductivity measurements because:

A. It increases temperature
B. It prevents electrolysis
C. It decreases ion movement
D. It increases resistance

Answer: B

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25. Primary batteries are:

A. Rechargeable
B. Non-rechargeable
C. Fuel cells
D. Electrolytic cells

Answer: B

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SECTION B: FILL IN THE BLANKS

1. The electrode where oxidation occurs is called ______.

Answer: Anode

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2. The electrode where reduction occurs is called ______.

Answer: Cathode

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3. Electrons flow from ______ to ______.

Answer: Anode to cathode

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4. Standard hydrogen electrode has electrode potential equal to ______.

Answer: 0.00 V

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5. The reciprocal of resistance is called ______.

Answer: Conductance

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6. The unit of resistance is ______.

Answer: Ohm (Ω)

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7. Kohlrausch’s law is applicable at ______ dilution.

Answer: Infinite

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8. The value of Faraday constant is approximately ______ C mol⁻¹.

Answer: 96500

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9. The process of coating a metal with zinc is called ______.

Answer: Galvanisation

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10. The product formed during rusting of iron is ______.

Answer: Hydrated ferric oxide

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11. The electrolyte used in lead storage battery is ______.

Answer: Sulphuric acid

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12. Fuel cells convert ______ energy into electrical energy.

Answer: Chemical

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13. Molar conductivity is represented by the symbol ______.

Answer: Λₘ

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14. The unit of molar conductivity is ______.

Answer: S cm² mol⁻¹

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15. Salt bridge contains an inert electrolyte like ______.

Answer: KCl/KNO₃

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SECTION C: ASSERTION–REASON QUESTIONS

Q1.

Assertion: A galvanic cell produces electricity.
Reason: The reaction occurring in a galvanic cell is spontaneous.

A. Both true and reason explains assertion
B. Both true but reason does not explain
C. Assertion true, reason false
D. Assertion false, reason true

Answer: A

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Q2.

Assertion: Molar conductivity increases on dilution.
Reason: Ion mobility increases and interionic attraction decreases.

Answer: A

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Q3.

Assertion: Salt bridge is necessary in a galvanic cell.
Reason: It maintains electrical neutrality.

Answer: A

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Q4.

Assertion: Primary batteries cannot be recharged.
Reason: Their chemical reactions are irreversible.

Answer: A

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Q5.

Assertion: Conductivity decreases on dilution.
Reason: Number of ions per unit volume decreases.

Answer: A

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SECTION D: VERY SHORT ANSWER QUESTIONS

1. What is electrochemistry?

Answer:
The branch of chemistry dealing with conversion between chemical energy and electrical energy.

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2. What is a salt bridge?

Answer:
A U-shaped tube containing electrolyte that connects two half cells.

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3. Define electrode potential.

Answer:
The potential difference developed between an electrode and electrolyte solution.

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4. Write the formula for Gibbs energy.

Answer:$$\Delta G=-nFE$$ΔG=−nFE

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5. State Faraday’s first law.

Answer:
Mass deposited is directly proportional to quantity of electricity passed.

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6. Name the rechargeable battery used in automobiles.

Answer:
Lead storage battery.

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7. Name the fuel used in hydrogen fuel cell.

Answer:
Hydrogen.

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8. What is corrosion?

Answer:
Slow deterioration of metals due to chemical or electrochemical reactions.

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9. What is the SI unit of conductivity?

Answer:
S m⁻¹

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10. What is the role of platinum in SHE?

Answer:
It acts as an inert electrode.

SECTION A: CASE-BASED QUESTIONS

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CASE STUDY 1: DANIELL CELL

A Daniell cell consists of a zinc electrode dipped in ZnSO₄ solution and a copper electrode dipped in CuSO₄ solution. The two solutions are connected through a salt bridge. Zinc undergoes oxidation while copper ions undergo reduction.

Cell representation:$$Zn|Zn^{2+}||Cu^{2+}|Cu$$Zn∣Zn2+∣∣Cu2+∣Cu

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Q1. The electrode where zinc oxidation occurs is:

A. Cathode
B. Anode
C. Salt bridge
D. Electrolyte

Answer: B

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Q2. The reaction at cathode is:

A.$$Zn\rightarrow Zn^{2+}+2e^-$$Zn→Zn2++2e−

B.$$Cu^{2+}+2e^-\rightarrow Cu$$Cu2++2e−→Cu

C.$$Cu\rightarrow Cu^{2+}+2e^-$$Cu→Cu2++2e−

D.$$Zn^{2+}+2e^-\rightarrow Zn$$Zn2++2e−→Zn

Answer: B

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Q3. The function of salt bridge is:

A. Supply electrons
B. Increase voltage
C. Maintain electrical neutrality
D. Stop ion movement

Answer: C

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Q4. Electrons move through external wire from:

A. Copper to zinc
B. Zinc to copper
C. Salt bridge to electrode
D. Cathode to anode

Answer: B

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CASE STUDY 2: CONDUCTIVITY OF ELECTROLYTES

Conductivity of an electrolyte solution depends on the number of ions present and their mobility. On dilution, conductivity decreases but molar conductivity increases.

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Q1. Conductivity is represented by:

A. R
B. κ
C. Λₘ
D. E°

Answer: B

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Q2. Molar conductivity increases on dilution because:

A. Number of ions decreases
B. Ionisation and mobility increase
C. Resistance becomes infinite
D. Temperature decreases

Answer: B

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Q3. Weak electrolytes show a sharp increase in molar conductivity because:

A. They become insoluble
B. Their ionisation increases
C. They stop conducting
D. Their mass increases

Answer: B

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Q4. Strong electrolytes are:

A. Completely ionised
B. Not ionised
C. Insoluble
D. Non-conducting

Answer: A

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CASE STUDY 3: ELECTROLYSIS

During electrolysis, electrical energy is converted into chemical energy. Reduction takes place at cathode and oxidation takes place at anode.

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Q1. In an electrolytic cell, cathode is:

A. Positive
B. Negative
C. Neutral
D. Absent

Answer: B

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Q2. Faraday’s first law states that:

A. Voltage depends on resistance
B. Mass deposited depends on charge passed
C. Current depends on temperature
D. Resistance depends on length

Answer: B

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Q3. Formula for deposited mass is:

A.$$m=IR$$m=IR

B.$$m=\frac{MIt}{nF}$$m=nFMIt​

C.$$m=VIt$$m=VIt

D.$$m=\frac{nF}{MI}$$m=MInF​

Answer: B

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SECTION B: ASSERTION–REASON QUESTIONS

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Q1.

Assertion: A positive value of E°cell indicates a spontaneous reaction.
Reason: ΔG° is negative for a spontaneous reaction.

Answer: A
(Both statements are true and reason correctly explains assertion.)

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Q2.

Assertion: Conductivity decreases on dilution.
Reason: Number of ions per unit volume decreases.

Answer: A

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Q3.

Assertion: Molar conductivity of weak electrolytes increases greatly on dilution.
Reason: Degree of ionisation increases on dilution.

Answer: A

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Q4.

Assertion: Salt bridge prevents direct mixing of solutions.
Reason: It contains an electrolyte in gel form.

Answer: A

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Q5.

Assertion: Lead storage battery can be recharged.
Reason: Its chemical reactions are reversible.

Answer: A

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SECTION C: MATCH THE FOLLOWING

Q1.

Column I	Column II
Anode	Oxidation
Cathode	Reduction
Salt bridge	Maintains neutrality
SHE	0 V

Answer:

• Anode → Oxidation
• Cathode → Reduction
• Salt bridge → Maintains neutrality
• SHE → 0 V

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Q2.

Column I	Column II
Dry cell	Primary battery
Lead storage battery	Secondary battery
Hydrogen fuel cell	Water as product
Rust	Fe₂O₃·xH₂O

Answer:

• Dry cell → Primary battery
• Lead storage battery → Secondary battery
• Hydrogen fuel cell → Water as product
• Rust → Fe₂O₃·xH₂O

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Q3.

Column I	Column II
Conductance	1/R
Cell constant	l/A
Gibbs energy	-nFE
Nernst equation	Non-standard potential

Answer:

• Conductance → 1/R
• Cell constant → l/A
• Gibbs energy → -nFE
• Nernst equation → Non-standard potential

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SECTION D: ONE-MARK IMPORTANT QUESTIONS

1. Which electrode is positive in a galvanic cell?

Answer: Cathode

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2. Which electrode is positive in an electrolytic cell?

Answer: Anode

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3. Name the law used to calculate limiting molar conductivity.

Answer: Kohlrausch’s law

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4. What is the value of n in Zn → Zn²⁺ + 2e⁻?

Answer: 2

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5. Which battery is used in automobiles?

Answer: Lead storage battery

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6. What is the product of hydrogen-oxygen fuel cell?

Answer: Water

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7. What type of reaction occurs at anode?

Answer: Oxidation

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8. What is the SI unit of molar conductivity?

Answer: S m² mol⁻¹

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9. Which gas is released at anode during electrolysis of molten NaCl?

Answer: Chlorine

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10. What protects iron from corrosion by zinc coating?

Answer: Galvanisation

PART 10: IMPORTANT LONG ANSWER QUESTIONS (3 & 5 MARKS)

Board Exam Writing Practice

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Q1. Explain the construction and working of a Daniell cell.

Answer:

A Daniell cell is a galvanic cell that converts chemical energy into electrical energy.

Construction:

It consists of:

• A zinc electrode dipped in ZnSO₄ solution.
• A copper electrode dipped in CuSO₄ solution.
• The two solutions are connected through a salt bridge.
• Electrodes are connected externally by a wire.

Cell notation:$$Zn|Zn^{2+}||Cu^{2+}|Cu$$Zn∣Zn2+∣∣Cu2+∣Cu

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Working:

At zinc electrode, oxidation occurs:$$Zn\rightarrow Zn^{2+}+2e^-$$Zn→Zn2++2e−

Zinc acts as anode.

At copper electrode, reduction occurs:$$Cu^{2+}+2e^-\rightarrow Cu$$Cu2++2e−→Cu

Copper acts as cathode.

Overall reaction:$$Zn+Cu^{2+}\rightarrow Zn^{2+}+Cu$$Zn+Cu2+→Zn2++Cu

Electrons flow from zinc electrode to copper electrode through the external circuit.

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Role of Salt Bridge:

1. Maintains electrical neutrality.
2. Completes the electrical circuit.
3. Prevents direct mixing of solutions.
4. Reduces liquid junction potential.

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Q2. Derive the Nernst equation for electrode potential.

Answer:

For an electrochemical reaction:$$Oxidised\ form+ne^-\rightarrow Reduced\ form$$Oxidised form+ne−→Reduced form

The Gibbs energy change is related to electrode potential:$$\Delta G=-nFE$$ΔG=−nFE

Under standard conditions:$$\Delta G^\circ=-nFE^\circ$$ΔG∘=−nFE∘

From thermodynamics:$$\Delta G=\Delta G^\circ+RT\ln Q$$ΔG=ΔG∘+RTlnQ

Substituting:$$-nFE=-nFE^\circ+RT\ln Q$$−nFE=−nFE∘+RTlnQ

Rearranging:$$E=E^\circ-\frac{RT}{nF}\ln Q$$E=E∘−nFRT​lnQ

This is the Nernst equation.

At 298 K:$$\boxed{E=E^\circ-\frac{0.0591}{n}\log Q}$$E=E∘−n0.0591​logQ​

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Q3. Derive the relation between Gibbs energy and cell potential.

Answer:

The maximum electrical work obtained from a cell is:$$W_{electrical}=nFE_{cell}$$Welectrical​=nFEcell​

Gibbs energy change is equal to negative of maximum useful work:$$\Delta G=-nFE_{cell}$$ΔG=−nFEcell​

For standard conditions:$$\boxed{\Delta G^\circ=-nFE^\circ_{cell}}$$ΔG∘=−nFEcell∘​​

Conclusions:

• If E°cell is positive → ΔG° is negative → reaction is spontaneous.
• If E°cell is negative → ΔG° is positive → reaction is non-spontaneous.

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Q4. Explain Kohlrausch’s law and its applications.

Answer:

Kohlrausch’s Law:

At infinite dilution, each ion contributes independently to the total molar conductivity of an electrolyte.

For an electrolyte:$$A_xB_y\rightarrow xA^{+}+yB^{-}$$Ax​By​→xA++yB− $$\Lambda_m^\circ=x\lambda_A^\circ+y\lambda_B^\circ$$Λm∘​=xλA∘​+yλB∘​

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Applications:

1. Finding limiting molar conductivity of weak electrolytes

Weak electrolytes cannot be measured directly at infinite dilution.

Example:$$\Lambda_m^\circ(CH_3COOH)$$Λm∘​(CH3​COOH)

is calculated using strong electrolytes.

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2. Finding degree of dissociation

$$\alpha=\frac{\Lambda_m}{\Lambda_m^\circ}$$α=Λm∘​Λm​​

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3. Finding dissociation constant

$$K_a=\frac{c\alpha^2}{1-\alpha}$$Ka​=1−αcα2​

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4. Finding solubility of sparingly soluble salts

$$S=\frac{1000\kappa}{\Lambda_m^\circ}$$S=Λm∘​1000κ​

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Q5. Explain the variation of molar conductivity with concentration.

Answer:

Molar conductivity is represented by:$$\Lambda_m$$Λm​

It changes differently for strong and weak electrolytes.

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Strong Electrolytes:

Examples:

• NaCl
• KCl
• HCl

Characteristics:

• Completely ionised.
• Molar conductivity increases slowly on dilution.

Reason:

• Interionic attraction decreases.
• Ion mobility increases.

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Weak Electrolytes:

Examples:

• CH₃COOH
• NH₄OH

Characteristics:

• Partially ionised.
• Molar conductivity increases sharply on dilution.

Reason:

• Degree of ionisation increases.
• More ions are produced.

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Q6. State and explain Faraday’s laws of electrolysis.

Answer:

Faraday’s First Law:

The mass of substance deposited at an electrode is directly proportional to the quantity of electricity passed.$$m\propto Q$$m∝Q

Since:$$Q=It$$Q=It

Therefore:$$m=ZIt$$m=ZIt

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Faraday’s Second Law:

When the same quantity of electricity is passed through different electrolytes, the masses deposited are proportional to their equivalent masses.$$m\propto \frac{M}{n}$$m∝nM​

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Combined equation:$$\boxed{m=\frac{MIt}{nF}}$$m=nFMIt​​

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Q7. Explain the working of a hydrogen-oxygen fuel cell.

Answer:

A hydrogen-oxygen fuel cell converts chemical energy directly into electrical energy.

Fuel:

• Hydrogen
• Oxygen

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Anode Reaction:

$$2H_2+4OH^-\rightarrow4H_2O+4e^-$$2H2​+4OH−→4H2​O+4e−

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Cathode Reaction:

$$O_2+2H_2O+4e^-\rightarrow4OH^-$$O2​+2H2​O+4e−→4OH−

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Overall Reaction:

$$2H_2+O_2\rightarrow2H_2O$$2H2​+O2​→2H2​O

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Advantages:

1. High efficiency.
2. Environment friendly.
3. Water is the only product.
4. Continuous electricity production is possible.

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Q8. Explain corrosion of iron and methods of prevention.

Answer:

Corrosion is the slow deterioration of metals due to chemical or electrochemical reactions with the environment.

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Rusting of Iron:

Anode reaction:

$$Fe\rightarrow Fe^{2+}+2e^-$$Fe→Fe2++2e−

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Cathode reaction:

$$O_2+2H_2O+4e^-\rightarrow4OH^-$$O2​+2H2​O+4e−→4OH−

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Formation of rust:$$Fe_2O_3.xH_2O$$Fe2​O3​.xH2​O

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Prevention Methods:

1. Painting

Prevents contact with air and moisture.

2. Galvanisation

Coating iron with zinc.

3. Electroplating

Depositing a protective metal layer.

4. Alloy formation

Example: Stainless steel.

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MOST IMPORTANT 5-MARK QUESTIONS TO PRACTICE

⭐ Derivation of Nernst equation
⭐ Daniell cell construction and working
⭐ Kohlrausch’s law and applications
⭐ Faraday laws with formula
⭐ Fuel cell working
⭐ Corrosion mechanism and prevention
⭐ Relation between ΔG°, E°cell and K

PART 11: ONE-SHOT REVISION SHEET + MIND MAP

ELECTROCHEMISTRY: COMPLETE CHAPTER MAP

                    ELECTROCHEMISTRY

                           |
        -----------------------------------------
        |                                       |
 Electrochemical Cells                 Electrolytic Cells
        |                                       |
  Chemical → Electrical              Electrical → Chemical
        |
  ----------------
  |              |
Anode          Cathode
Oxidation     Reduction

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1. ELECTROCHEMICAL CELLS

Galvanic Cell

Definition:

A cell that converts chemical energy into electrical energy through a spontaneous reaction.

Example:

Daniell Cell:$$Zn|Zn^{2+}||Cu^{2+}|Cu$$Zn∣Zn2+∣∣Cu2+∣Cu

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Electrode Reactions:

Anode:

$$Zn\rightarrow Zn^{2+}+2e^-$$Zn→Zn2++2e−

(Oxidation)

Cathode:

$$Cu^{2+}+2e^-\rightarrow Cu$$Cu2++2e−→Cu

(Reduction)

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Remember:

AN OX

Anode → Oxidation

RED CAT

Reduction → Cathode

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2. SALT BRIDGE

Functions:

✔ Maintains electrical neutrality
✔ Completes electrical circuit
✔ Prevents direct mixing of solutions
✔ Minimises liquid junction potential

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3. ELECTRODE POTENTIAL

Standard Conditions:

• Temperature = 298 K
• Concentration = 1 M
• Pressure = 1 bar

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Standard Hydrogen Electrode (SHE)

$$E^\circ=0.00V$$E∘=0.00V

Used as reference electrode.

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4. CELL EMF

Formula:$$\boxed{E^\circ_{cell}=E^\circ_{cathode}-E^\circ_{anode}}$$Ecell∘​=Ecathode∘​−Eanode∘​​

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Spontaneity Rule:

Positive E°cell

→ Spontaneous reaction

→ ΔG° negative

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Negative E°cell

→ Non-spontaneous reaction

→ ΔG° positive

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5. NERNST EQUATION

Used for calculating electrode potential under non-standard conditions.$$\boxed{E=E^\circ-\frac{0.0591}{n}\log Q}$$E=E∘−n0.0591​logQ​

(at 298 K)

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Reaction Quotient:

Products concentration / Reactants concentration

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6. IMPORTANT RELATIONS

Gibbs Energy:

$$\boxed{\Delta G=-nFE}$$ΔG=−nFE​

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Standard Gibbs Energy:

$$\boxed{\Delta G^\circ=-nFE^\circ}$$ΔG∘=−nFE∘​

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Equilibrium Constant:

$$\boxed{E^\circ=\frac{0.0591}{n}\log K}$$E∘=n0.0591​logK​

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7. CONDUCTANCE

Resistance

$$R=\rho\frac{l}{A}$$R=ρAl​

Unit:

Ω

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Conductance

$$G=\frac1R$$G=R1​

Unit:

Siemens (S)

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Conductivity

$$\kappa=\frac1\rho$$κ=ρ1​

Unit:

S m⁻¹

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8. CELL CONSTANT

Formula:$$\boxed{\frac{l}{A}}$$Al​​

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Conductivity:$$\boxed{\kappa=\frac{\text{Cell Constant}}{R}}$$κ=RCell Constant​​

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9. MOLAR CONDUCTIVITY

Formula:$$\boxed{\Lambda_m=\frac{\kappa\times1000}{M}}$$Λm​=Mκ×1000​​

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Effect of Dilution:

Conductivity:

Decreases

Reason:

Number of ions per unit volume decreases.

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Molar Conductivity:

Increases

Reason:

• Ion mobility increases.
• Interionic attraction decreases.

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10. STRONG VS WEAK ELECTROLYTES

Strong Electrolyte	Weak Electrolyte
Completely ionised	Partially ionised
High conductivity	Low conductivity
Small increase in Λm	Large increase in Λm
NaCl, HCl	CH₃COOH

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11. KOHLRAUSCH’S LAW

Statement:

At infinite dilution, each ion contributes independently to the total molar conductivity.

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Formula:$$\Lambda_m^\circ=\lambda^\circ_+ +\lambda^\circ_-$$Λm∘​=λ+∘​+λ−∘​

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Applications:

✔ Calculate Λ°m of weak electrolytes

✔ Calculate degree of dissociation$$\alpha=\frac{\Lambda_m}{\Lambda_m^\circ}$$α=Λm∘​Λm​​

✔ Calculate dissociation constant$$K_a=\frac{c\alpha^2}{1-\alpha}$$Ka​=1−αcα2​

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12. ELECTROLYSIS

Electrolytic Cell:

Electrical energy → Chemical energy

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Electrode Rules:

Cathode:

Negative

Reduction occurs

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Anode:

Positive

Oxidation occurs

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13. FARADAY’S LAWS

First Law:

Mass deposited is directly proportional to charge passed.$$m\propto Q$$m∝Q

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Second Law:

Mass deposited depends on equivalent mass.

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Formula:

$$\boxed{m=\frac{MIt}{nF}}$$m=nFMIt​​

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Faraday Constant:

$$\boxed{F=96500Cmol^{-1}}$$F=96500Cmol−1​

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14. BATTERIES

Primary Batteries

• Non-rechargeable
• Irreversible reactions

Examples:

✔ Dry cell
✔ Mercury cell

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Secondary Batteries

• Rechargeable
• Reversible reactions

Examples:

✔ Lead storage battery
✔ Lithium-ion battery

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15. FUEL CELLS

Hydrogen-Oxygen Fuel Cell

Reaction:$$2H_2+O_2\rightarrow2H_2O$$2H2​+O2​→2H2​O

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Advantages:

✔ High efficiency
✔ Environment friendly
✔ Water is the only product

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16. CORROSION

Definition:

Slow destruction of metals due to chemical/electrochemical reactions.

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Rust Formation

Anode:

$$Fe\rightarrow Fe^{2+}+2e^-$$Fe→Fe2++2e−

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Cathode:

$$O_2+2H_2O+4e^-\rightarrow4OH^-$$O2​+2H2​O+4e−→4OH−

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Rust:$$Fe_2O_3.xH_2O$$Fe2​O3​.xH2​O

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Prevention:

✔ Painting
✔ Oiling
✔ Galvanisation
✔ Electroplating
✔ Alloy formation

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LAST 15-MINUTE REVISION TABLE

Topic	Must Remember
SHE	0 V
Electron flow	Anode → Cathode
Oxidation	Anode
Reduction	Cathode
E°cell formula	Cathode − Anode
ΔG formula	-nFE
Nernst equation	E = E° − 0.0591/n log Q
Faraday constant	96500 C mol⁻¹
Salt bridge	Maintains neutrality
Rust	Fe₂O₃·xH₂O
Fuel cell product	Water