ELECTROCHEMISTRY
CHAPTER - 03
3.1 How would you determine the standard electrode potential of the system Mg2+|Mg?
3.2 Can you store copper sulphate solutions in a zinc pot?
Ans. Zinc is more reactive than copper. Therefore, zinc can displace copper from its salt solution. If copper sulphate solution is stored in a zinc pot, then zinc will displace copper from the copper sulphate solution.
Zn+CuSO4 ⟶ZnSO4 +Cu
Hence, the copper sulphate solution cannot be stored in a zinc pot.
3.3 Consult the table of standard electrode potentials and suggest three substances that can oxidise ferrous ions under suitable conditions.
Ans. The metals having higher oxidation potential than iron are able to oxidize iron.
∴Ag+ + e- ⟶Ag E0 = 0.80 V
Br2 + 2e- ⟶ Br- E0 = 1.06 V
Cl2 + 2e- ⟶ Cl⊝ E0 =1.36 V
Example 3.1 Represent the cell in which the following reaction takes place
Mg(s) + 2Ag+ (0.0001M) → Mg2+(0.130M) + 2Ag(s)
Calculate its Ecell if E0cell = 3.17 V.
Ans.
Example 3.2 Calculate the equilibrium constant of the reaction:
Cu(s) + 2Ag+ (aq) → Cu2+(aq) + 2Ag(s)
E0cell = 0.46 V
Solution: E0cell = 0.059 V /2 log Kc = 0.46 V
log Kc = 0.46 V x 2 / 0.059 V = 15.6
Kc = 3.92 x 1015
Example 3.3 The standard electrode potential for Daniell cell is 1.1 V. Calculate the standard Gibbs energy for the reaction.
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu (s)
Solution: ΔrG0 = -nFE0cell
n in the above equation is 2.
F = 96487 C /mol and E0cell = 1.1V
Therefore
ΔrG0 = -2 x 1.1 V x 96487 C /mol = -21227 J / mol = -212.27 kJ/ mol
Intext Questions: 3.4 Calculate the potential of hydrogen electrode in contact with a solution whose pH is 10.
Intext Questions: 3.5 Calculate the emf of the cell in which the following reaction takes place:
Ni(s) + 2Ag+ (0.002 M) → Ni2+ (0.160 M) + 2Ag(s)
Given that E0cell = 1.05 V
Intext Questions: 3.6 The cell in which the following reaction occurs:
2Fe3+ (aq) + 2I- (aq) → 2Fe2+(aq) + I2 (s) has E0cell at 298K.
Calculate the standard Gibbs energy and the equilibrium constant of the cell reaction.
Example 3.4 Resistance of a conductivity cell filled with 0.1 mol L–1 KCl solution is 100 Ω. If the resistance of the same cell when filled with 0.02 mol L–1 KCl solution is 520 Ω, calculate the conductivity and molar conductivity of 0.02 mol L–1 KCl solution. The conductivity of 0.1 mol L–1 KCl solution is 1.29 S/m
Ans.
Example 3.5 The electrical resistance of a column of 0.05 mol L–1 NaOH solution of diameter 1 cm and length 50 cm is 5.55 × 103 ohm. Calculate its resistivity, conductivity and molar conductivity.
Intext Question 3.7 Why does the conductivity of a solution decrease with dilution?
Ans. Conductivity of a solution is the conductance of ions present in a unit volume of the solution. On dilution, the number of ions per unit volume decreases. Hence the conductivity decreases.
Intext Question 3.8 3.8 Suggest a way to determine the Λ° m value of water.
Intext Questions 3.9 The molar conductivity of 0.025 mol L–1 methanoic acid is 46.1 S cm2 mol–1 . Calculate its degree of dissociation and dissociation constant. Given λ 0 (H+ ) = 349.6 S cm2 mol–1 and λ 0 (HCOO– ) = 54.6 S cm2 mol–1 .
Example 3.10 A solution of CuSO4 is electrolysed for 10 minutes with a current of 1.5 amperes. What is the mass of copper deposited at the cathode?
t = 600 s charge = current × time = 1.5 A × 600 s = 900 C
According to the reaction: Cu2+(aq) + 2e– = Cu(s)
We require 2F or 2 × 96487 C to deposit 1 mol or 63 g of Cu.
For 900 C, the mass of Cu deposited = (63 g mol–1 × 900 C)/(2 × 96487 C mol–1) = 0.2938 g.
Intext Question 3.10 If a current of 0.5 ampere flows through a metallic wire for 2 hours, then how many electrons would flow through the wire?
Intext Question 3.11 Suggest a list of metals that are extracted electrolytically.
Ans. Metals that are on the top of the reactivity series such as sodium, potassium, calcium, lithium, magnesium, aluminium are extracted electrolytically.
Intext Question 3.12 Consider the reaction: Cr2O7 2– + 14H+ + 6e– → 2Cr3+ + 7H2O
What is the quantity of electricity in coulombs needed to reduce 1 mol of Cr2O7 2–?
