Chapter 2: Electrochemistry Short Questions | chemistry Class 12 CBSE

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Chapter 2: Electrochemistry Short Questions | chemistry Class 12 CBSE | ByteNotes.in

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Question 1

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Distinguish between a galvanic cell and an electrolytic cell.

A galvanic cell converts chemical energy of a spontaneous redox reaction into electrical energy, whereas an electrolytic cell uses electrical energy to carry out a non-spontaneous chemical reaction.
In a galvanic cell the anode is negative and cathode is positive; in an electrolytic cell the anode is connected to the positive terminal and cathode to the negative terminal of an external source.
Examples: Daniell cell is galvanic; electrolysis of molten NaCl is electrolytic.

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Question 2

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What is meant by standard electrode potential? How is it measured?

Standard electrode potential is the electrode potential when the concentration of all species involved in the half-cell is unity (1 mol dm−3), measured against the Standard Hydrogen Electrode.
It is measured by constructing a cell with the SHE as the reference anode and the given electrode as cathode; the measured EMF equals E°R since E°L for SHE is zero.
IUPAC convention calls these standard reduction potentials (standard electrode potentials), and the sign indicates whether the reduced form is more or less stable than H2.

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Question 3

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State the relationship between standard Gibbs energy and standard cell potential.

The relationship is ΔrG° = −nFE°(cell), where n is the number of moles of electrons transferred and F is the Faraday constant (96487 C mol−1).
E°(cell) is an intensive property (independent of stoichiometric coefficients), but ΔrG° is extensive and changes when the equation is multiplied.
This equation enables calculation of thermodynamic quantities like ΔrG° and hence equilibrium constant K (via ΔrG° = −RT ln K) from measurable cell potentials.

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Question 4

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Write the Nernst equation for a cell reaction and explain its significance.

For the general reaction aA + bB → cC + dD, the Nernst equation is: E(cell) = E°(cell) − (RT/nF) ln([C]c[D]d / [A]a[B]b).
At 298 K, substituting values of R, T and F, it simplifies to E(cell) = E°(cell) − (0.059/n) log Q.
Its significance is that it shows cell potential depends on concentration of reactants and products; at equilibrium E(cell) = 0, yielding E°(cell) = (0.059/n) log Kc.

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Question 5

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Define resistivity and conductivity. State their SI units and interrelation.

Resistivity (ρ) is the resistance of a substance when it is one metre long with one square metre cross-section; SI unit is ohm metre (Ω m).
Conductivity (κ) is the inverse of resistivity (κ = 1/ρ) and represents the ability of a material to conduct electricity; SI unit is S m−1.
They are related by κ = 1/ρ; for a conductor of length l and area A: R = ρl/A and G = κA/l.

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Question 6

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Why is an AC source used instead of a DC source for measuring resistance of ionic solutions?

Passing direct current (DC) through an ionic solution over a prolonged period leads to change in composition due to electrochemical reactions (electrolysis).
Alternating current (AC) avoids this problem as it reverses direction rapidly, preventing sustained electrode reactions and compositional changes.
An AC source in the audio frequency range (550–5000 cycles per second) fed through a Wheatstone bridge setup allows accurate resistance measurement without altering the solution.

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Question 7

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Define molar conductivity. How does it differ from conductivity?

Molar conductivity (Λm) is defined as the conductivity of a solution divided by its molar concentration: Λm = κ/c, with units S cm2 mol−1 or S m2 mol−1.
Conductivity (κ) measures the conductance per unit volume and decreases with dilution because the number of ions per unit volume decreases.
Molar conductivity increases with dilution because the total volume containing one mole of electrolyte increases, more than compensating for the decrease in κ.

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Question 8

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State Kohlrausch's law of independent migration of ions.

Kohlrausch's law states that the limiting molar conductivity of an electrolyte can be represented as the sum of the individual contributions of the anion and cation of the electrolyte.
Mathematically: Λ°m = ν+λ°+ + ν−λ°−, where ν+ and ν− are the number of cations and anions per formula unit, and λ°+ and λ°− are their respective limiting molar conductivities.
For example, Λ°m(NaCl) = λ°Na+ + λ°Cl−. This law enables calculation of Λ°m for weak electrolytes and determination of dissociation constants.

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Question 9

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State Faraday's first and second laws of electrolysis.

Faraday's First Law: The amount of chemical reaction at any electrode during electrolysis is proportional to the quantity of electricity passed through the electrolyte (solution or melt).
Faraday's Second Law: The amounts of different substances liberated by the same quantity of electricity are proportional to their chemical equivalent weights (atomic mass ÷ number of electrons required to reduce the cation).
One Faraday (≈ 96500 C mol−1) is the charge on one mole of electrons and deposits one equivalent of substance.

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Question 10

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Describe the construction of the Standard Hydrogen Electrode (SHE).

SHE consists of a platinum electrode coated with platinum black, dipped in an acidic solution of 1 M H+ concentration, with pure hydrogen gas bubbled through it at 1 bar pressure.
The platinum black increases the surface area for adsorption of H2 and acts as an inert conductor. Concentrations of both reduced (H2) and oxidised (H+) forms are maintained at unity.
SHE is assigned a potential of zero at all temperatures; the half-reaction is H+(aq) + e− → ½H2(g). It serves as the universal reference for measuring electrode potentials.

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Question 11

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How does the conductivity of ionic solutions depend on the nature and concentration of the electrolyte?

Conductivity depends on the nature of the electrolyte — strong electrolytes (fully dissociated) give higher conductivity than weak electrolytes at the same concentration.
It depends on ion size and solvation: smaller, less solvated ions move faster under a potential gradient, giving higher conductivity.
Conductivity always decreases with decreasing concentration because the number of ions per unit volume carrying current decreases on dilution.

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Question 12

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What is the role of a salt bridge in a galvanic cell?

A salt bridge provides an ionic connection between the two half-cells, allowing migration of ions to maintain electrical neutrality in each half-cell compartment.
Without a salt bridge, charge would build up in each compartment (excess positive charge in one and excess negative charge in the other), rapidly stopping current flow.
It typically contains a concentrated solution of an inert electrolyte (like KCl or KNO3) in a gel, ensuring that neither the cation nor anion participates in the cell reactions.

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Question 13

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What is meant by cell constant? How is it determined experimentally?

Cell constant (G*) is defined as the ratio l/A, where l is the distance between the electrodes and A is their area of cross-section; it has the dimensions of length−1 (cm−1 or m−1).
Since direct measurement of l and A is unreliable, G* is determined by measuring the resistance of a KCl solution of known conductivity: G* = R × κ.
Once determined, the cell constant allows calculation of conductivity for any unknown solution as κ = G*/R.

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Chapter 2: Electrochemistry | Short Questions

Distinguish between a galvanic cell and an electrolytic cell.

What is meant by standard electrode potential? How is it measured?

State the relationship between standard Gibbs energy and standard cell potential.

Write the Nernst equation for a cell reaction and explain its significance.

Define resistivity and conductivity. State their SI units and interrelation.

Why is an AC source used instead of a DC source for measuring resistance of ionic solutions?

Define molar conductivity. How does it differ from conductivity?

State Kohlrausch's law of independent migration of ions.

State Faraday's first and second laws of electrolysis.

Describe the construction of the Standard Hydrogen Electrode (SHE).

How does the conductivity of ionic solutions depend on the nature and concentration of the electrolyte?

What is the role of a salt bridge in a galvanic cell?

What is meant by cell constant? How is it determined experimentally?

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