JEE Chemistry Electrochemistry MCQs Set 03

Answer: (c) Dissociate into ions even at high concentration

Explanation:
1. A strong electrolyte is almost completely ionised at all concentrations.
2. Solubility and the ability to conduct do not define strength.

Teacher's Note:
1. Strong electrolytes: strong acids, strong bases and most salts.
2. Weak electrolytes ionise only slightly, more at dilution.

Answer: (b) Free ions

Explanation:
1. In the molten state the lattice breaks into mobile Na+ and Cl- ions.
2. Ionic conduction needs no free electrons.

Teacher's Note:
1. Ionic solids conduct only when molten or in solution.
2. Metals conduct by electrons.

Answer: (c) Ions

Explanation:
1. In an electrolyte the current is carried by cations and anions.
2. Electrons carry the current in metals.

Teacher's Note:
1. Electrolytic conduction is ionic.
2. It is accompanied by chemical change at electrodes.

Answer: (a) 1 M

Explanation:
1. Equivalent conductance increases with dilution because ionic mobility increases and ionisation rises.
2. The most concentrated solution has the lowest value.

Teacher's Note:
1. Lambda increases on dilution for all electrolytes.
2. It is highest at infinite dilution.

Answer: (d) 10,000 cc

Explanation:
1. A 1 N solution has 1 equivalent in 1000 cc.
2. For N/10, 1 equivalent is present in 10 times this volume, 10,000 cc.

Teacher's Note:
1. V = 1000 / normality.
2. More dilute means larger V.

Answer: (c) \( \dfrac{1000}{RC} \)

Explanation:
1. \( \kappa=1/R \) and \( \Lambda=\kappa\times1000/C \).
2. So \( \Lambda=1000/(RC) \).

Teacher's Note:
1. Specific conductance is the reciprocal of specific resistance.
2. 1000 converts litres into cc.

Answer: (c) R

Explanation:
1. Resistance is inversely proportional to conductivity.
2. R has the smallest conductivity, so it has the highest resistance.

Teacher's Note:
1. Lowest conductance means highest resistance.
2. Compare the powers of ten carefully.

Answer: (c) 10.61

Explanation:
1. \( \Lambda_m=\kappa\times1000/M \).
2. \( =1.061\times10^{-4}\times1000/0.01=10.61 \) S cm\(^2\) mol\(^{-1}\).

Teacher's Note:
1. Multiply by 1000 and divide by molarity.
2. Units are S cm^2 mol^-1.

Answer: (d) 1.0 N

Explanation:
1. Specific conductance is the conductance of 1 cc and depends on the number of ions per cc.
2. It increases with concentration.

Teacher's Note:
1. Specific conductance rises with concentration; equivalent conductance falls.
2. Dilution lowers kappa.

Answer: (b) The velocity of H+ ions is more than that of Na+ ions

Explanation:
1. H+ ions move very fast by the Grotthuss (proton jumping) mechanism.
2. Both HCl and NaCl are strong electrolytes, so ionisation is not the reason.

Teacher's Note:
1. H+ and OH- have the highest ionic mobilities.
2. Mobility, not mass, decides conductivity.

Answer: (d) \( 4\times10^{-6} \)

Explanation:
1. Solubility \( S=\kappa\times1000/\Lambda_m=3.06\times10^{-6}\times1000/1.53=2\times10^{-3} \) mol/L.
2. \( K_{sp}=S^2=4\times10^{-6} \).

Teacher's Note:
1. For sparingly soluble salts Lambda is taken at infinite dilution.
2. Ksp = s^2 for a 1:1 salt.

Answer: (b) From Cu to Ag

Explanation:
1. Cu is the anode, where oxidation gives electrons, and Ag is the cathode.
2. Electrons flow from anode to cathode, from Cu to Ag.

Teacher's Note:
1. Anode to cathode through the wire.
2. E(Ag) is higher than E(Cu).

Answer: (d) \( \ce{Zn^{2+}(aq) + 2e- -> Zn(s)} \)

Explanation:
1. Standard reduction potential refers to the reduction of the aqueous ion to the metal.
2. Gas-phase ions or the oxidation direction do not define it.

Teacher's Note:
1. State symbols matter: ion in aqueous solution.
2. The metal is a solid.

Answer: (b) The electrode potential is assumed to be zero

Explanation:
1. The absolute potential of a single electrode cannot be measured.
2. By convention the potential of the SHE is taken to be zero.

Teacher's Note:
1. SHE: Pt, H2 (1 bar)
2. H+ (1 M) at 298 K.
3. All other potentials are relative to it.

Answer: (d) It contains a metallic conductor which does not adsorb hydrogen

Explanation:
1. The platinum used is coated with platinum black, which adsorbs hydrogen strongly.
2. The other three conditions are part of the definition.

Teacher's Note:
1. Platinum black adsorbs H2 gas.
2. That makes the electrode reversible.

Answer: (b) +0.76 V

Explanation:
1. \( E^\circ_{cell}=E^\circ_{cathode}-E^\circ_{anode}=0-(-0.76) \).
2. \( E^\circ_{cell}=+0.76 \) V.

Teacher's Note:
1. Cathode minus anode, both as reduction potentials.
2. Positive emf means spontaneous.

Answer: (c) Pt | H\(_2\)(g), 1 bar | 1 M HCl(aq) | AgCl(s) | Ag(s)

Explanation:
1. Hydrogen is oxidised at the anode (Pt, H2) and AgCl is reduced at the Ag-AgCl electrode.
2. Both electrodes dip in HCl, so the electrolyte is HCl.

Teacher's Note:
1. The anode is written on the left.
2. Ag-AgCl is a metal-insoluble salt electrode.

Answer: (b) Negative

Explanation:
1. \( E^\circ=E^\circ_{Zn}-E^\circ_{Cu}=-0.76-0.34=-1.10 \) V.
2. The emf is negative, so the cell reaction is non spontaneous as written.

Teacher's Note:
1. The right electrode is the cathode.
2. Zinc is a stronger reductant than copper.

Answer: (c) \( 0.771-0.536=0.235 \) V

Explanation:
1. Fe3+ is reduced (cathode) and I- is oxidised (anode).
2. E = 0.771 - 0.536 = 0.235 V.
3. Electrode potentials are intensive and are not multiplied by coefficients.

Teacher's Note:
1. Never multiply E° by stoichiometric numbers.
2. Cathode minus anode.

Answer: (a) \( \ce{Ag+ -> Ag - e-} \) (that is, \( \ce{Ag+ + e- -> Ag} \))

Explanation:
1. The first reaction is the gain of an electron by Ag+, which is reduction, so it occurs at the cathode.
2. The others are oxidations (loss of electrons) and are anodic.

Teacher's Note:
1. Anode = oxidation.
2. Check whether electrons are on the left or right.