Kerala Plus Two Chemistry Electrochemistry Important Questions

 Electrochemistry can be challenging because it combines chemical reactions, electrical energy and numerical calculations. Concepts such as electrode potential, molar conductivity, the Nernst equation and Faraday’s laws require you to understand the underlying principles and apply the correct formulas to solve problems.

These Kerala Plus Two Chemistry Electrochemistry Important Questions cover key concepts from Chapter 2 through theory questions, numerical problems and MCQs. Practise topics such as electrochemical cells, Kohlrausch’s law, batteries, fuel cells and corrosion to strengthen your understanding and prepare for the Kerala DHSE Class 12 Chemistry examination.

Important Topics in Electrochemistry for Kerala Plus Two

Revise the following concepts while preparing for Kerala Plus Two Chemistry Chapter 2:

You can also refer to the Kerala Class 12 syllabus to check the prescribed topics. Reviewing the Kerala Plus Two exam pattern can help you understand the examination format and plan your revision.

Kerala Plus Two Chemistry Chapter 2 Important Questions 

Practising important questions can help you identify concepts that require further revision. Use the following Kerala DHSE Plus Two Chemistry Chapter 2 MIQs to review theory, calculations and objective questions.

Q1. State Kohlrausch’s law of independent migration of ions. Explain how it helps determine the limiting molar conductivity of a weak electrolyte such as acetic acid.

Answer:

Kohlrausch’s law states that at infinite dilution, each ion contributes independently to the total limiting molar conductivity of an electrolyte.

The limiting molar conductivity of acetic acid can be calculated using:

[
\begin{aligned}
\Lambda_m^\circ(\mathrm{CH_3COOH})={}&
\Lambda_m^\circ(\mathrm{HCl})\
&+\Lambda_m^\circ(\mathrm{CH_3COONa})\
&-\Lambda_m^\circ(\mathrm{NaCl})
\end{aligned}
]

This method helps determine the limiting molar conductivity of weak electrolytes.

Q2. The standard electrode potential of a Daniell cell is 1.1 V. Calculate the standard Gibbs energy change for the reaction:

[
\mathrm{Zn(s)+Cu^{2+}(aq)\rightarrow Zn^{2+}(aq)+Cu(s)}
]

Answer:

[
\Delta G^\circ=-nFE^\circ_{\mathrm{cell}}
]

Given: (n=2), (F=96500\ \mathrm{C,mol^{-1}}), and (E^\circ_{\mathrm{cell}}=1.1\ \mathrm{V}).

[
\begin{aligned}
\Delta G^\circ&=-2\times96500\times1.1\
&=-212300\ \mathrm{J,mol^{-1}}
\end{aligned}
]

[
\boxed{\Delta G^\circ=-212.3\ \mathrm{kJ,mol^{-1}}}
]

The negative value indicates that the reaction is spontaneous under standard conditions.

Q3. Explain the construction and working of the (\mathrm{H_2-O_2}) fuel cell, including the chemical equations for the electrode reactions.

Answer:

A hydrogen–oxygen fuel cell converts the chemical energy of hydrogen and oxygen directly into electrical energy. In an alkaline fuel cell, hydrogen is supplied to the anode, oxygen to the cathode, and aqueous KOH acts as the electrolyte.

Anode reaction:

[
2\mathrm{H_2}+4\mathrm{OH^-}\rightarrow4\mathrm{H_2O}+4e^-
]

Cathode reaction:

[
\mathrm{O_2}+2\mathrm{H_2O}+4e^-\rightarrow4\mathrm{OH^-}
]

Overall reaction:

[
2\mathrm{H_2}+\mathrm{O_2}\rightarrow2\mathrm{H_2O}
]

Electrons released at the anode flow through the external circuit to the cathode, generating electric current.

Q4. Write the anode and cathode reactions during the operation of a lead storage battery. Mention the electrolyte used.

Answer:

A lead storage battery is a rechargeable secondary cell. It uses lead ((\mathrm{Pb})) as the negative electrode, lead dioxide ((\mathrm{PbO_2})) as the positive electrode, and aqueous sulphuric acid ((\mathrm{H_2SO_4})) as the electrolyte.

During discharge:

Anode reaction:

[
\mathrm{Pb+SO_4^{2-}\rightarrow PbSO_4+2e^-}
]

Cathode reaction:

[
\mathrm{PbO_2+4H^++SO_4^{2-}+2e^-
\rightarrow PbSO_4+2H_2O}
]

Overall reaction:

[
\mathrm{Pb+PbO_2+2H_2SO_4
\rightarrow2PbSO_4+2H_2O}
]

Q5. Briefly explain the electrochemical processes involved in the rusting of iron.

Answer:

Rusting occurs when iron is exposed to moisture and oxygen. Different regions on the iron surface act as anode and cathode.

Anode reaction:

[
\mathrm{Fe\rightarrow Fe^{2+}+2e^-}
]

Cathode reaction:

[
\mathrm{O_2+2H_2O+4e^-\rightarrow4OH^-}
]

The iron ions react further with hydroxide ions and oxygen to form hydrated iron(III) oxide, commonly known as rust.

Q6. A solution of (\mathrm{CuSO_4}) is electrolysed for 20 minutes using a current of 1.5 A. Calculate the mass of copper deposited at the cathode. Given: atomic mass of Cu = 63.

Answer:

Using Faraday’s first law:

[
m=\frac{MIt}{nF}
]

Given: (M=63\ \mathrm{g,mol^{-1}}), (I=1.5\ \mathrm{A}), (t=1200\ \mathrm{s}), (n=2), and (F=96500\ \mathrm{C,mol^{-1}}).

[
m=\frac{63\times1.5\times1200}{2\times96500}
]

[
\boxed{m\approx0.587\ \mathrm{g}}
]

Therefore, approximately 0.587 g of copper is deposited at the cathode.

Q7. Define molar conductivity. How does it vary with concentration for strong and weak electrolytes?

Answer:

Molar conductivity ((\Lambda_m)) is the conductance of the volume of solution containing one mole of an electrolyte.

It is expressed as:

[
\Lambda_m=\frac{\kappa\times1000}{C}
]

where (\kappa) is conductivity in (\mathrm{S,cm^{-1}}) and (C) is concentration in (\mathrm{mol,L^{-1}}).

• Strong electrolytes: Molar conductivity increases gradually on dilution because interionic interactions decrease.

• Weak electrolytes: Molar conductivity increases sharply on dilution because ionisation increases significantly.

For both types, molar conductivity approaches its limiting value at infinite dilution.

Q8. Calculate the electrode potential of a hydrogen electrode in contact with a 0.01 M (\mathrm{H^+}) solution. Given (E^\circ_{\mathrm{H^+/H_2}}=0\ \mathrm{V}). Assume hydrogen gas is at 1 bar and the temperature is 298 K.

Answer:

The electrode reaction is:

[
2\mathrm{H^+}+2e^-\rightarrow\mathrm{H_2(g)}
]

Using the Nernst equation:

[
E=E^\circ-\frac{0.0591}{2}
\log\frac{P_{\mathrm{H_2}}}{[\mathrm{H^+}]^2}
]

Substituting (P_{\mathrm{H_2}}=1) bar and ([\mathrm{H^+}]=0.01\ \mathrm{M}):

[
E=0-\frac{0.0591}{2}\log\frac{1}{(0.01)^2}
]

[
\boxed{E=-0.0591\ \mathrm{V}}
]

Q9. Which of the following is an example of a secondary cell?

(A) Dry cell

(B) Lead storage battery

(C) Mercury cell

(D) Leclanché cell

Answer: (B) Lead storage battery

A lead storage battery is rechargeable, whereas the other listed cells are primary cells.

Q10. A galvanic cell is represented as:

[
\mathrm{Mg(s)|Mg^{2+}(aq)||Ag^+(aq)|Ag(s)}
]

Identify the anode and cathode, and write the Nernst equation for its cell potential.

Answer:

• Anode: Magnesium electrode, where oxidation occurs.

• Cathode: Silver electrode, where reduction occurs.

Anode reaction:

[
\mathrm{Mg\rightarrow Mg^{2+}+2e^-}
]

Cathode reaction:

[
\mathrm{2Ag^++2e^-\rightarrow2Ag}
]

Overall reaction:

[
\mathrm{Mg+2Ag^+\rightarrow Mg^{2+}+2Ag}
]

The Nernst equation at 298 K is:

[
E_{\mathrm{cell}}=E^\circ_{\mathrm{cell}}
-\frac{0.0591}{2}\log
\frac{[\mathrm{Mg^{2+}}]}{[\mathrm{Ag^+}]^2}
]

Pure solids are omitted from the reaction quotient because their activities are taken as unity.

Kerala Plus Two Chemistry Electrochemistry Important Questions PDF

Electrochemistry includes important concepts, formulas and numerical problems that can be challenging to revise before the Kerala Plus Two Chemistry examination. Practising key questions can help you understand these concepts and prepare for different question types.

How to Prepare With Electrochemistry Important Questions 

Revise key formulas, practise numerical problems, and review electrode reactions to prepare effectively with Electrochemistry Important Questions. Use these tips to revise Electrochemistry effectively:

• Memorise essential formulas: Revise the Nernst equation, the Gibbs energy relationship and Faraday’s laws.

• Practise numerical problems: Solve questions on cell potential, Gibbs energy and the mass of substances deposited during electrolysis.

• Learn electrode reactions: Write the anode and cathode reactions for galvanic cells, lead storage batteries and fuel cells.

• Understand conductivity: Know how conductivity and molar conductivity differ and how they change with dilution.

• Revise conceptual differences: Distinguish between primary and secondary cells, galvanic and electrolytic cells, and strong and weak electrolytes.

• Solve questions without checking answers first: Attempt each question independently before reviewing its solution.

• Plan your revision: Check the Kerala Plus Two timetable for 2027 for examination planning, using the relevant timetable for your academic session.

• Understand the marking scheme: Refer to the Kerala Plus Two grading and marking system to understand how marks are awarded.

Electrochemistry becomes easier to revise when you connect its formulas with the chemical processes they describe. Practise numerical questions on electrode potential, Gibbs energy and electrolysis alongside theory questions on conductivity, batteries, fuel cells and corrosion. Regular revision of these Kerala Plus Two Chemistry Electrochemistry Important Questions can help you identify gaps in your understanding and improve your accuracy when solving Chapter 2 questions.