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SS2: CHEMISTRY - 1ST TERM

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  1. Periodicity and Periodic Table I | Week 1
    5 Topics
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    1 Quiz
  2. Quantum Numbers Orbitals & Electrical Structure | Week 2
    6 Topics
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    1 Quiz
  3. Periodicity and Periodic Table II | Week 3
    12 Topics
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    1 Quiz
  4. Periodicity and Periodic Properties III | Week 4
    11 Topics
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    1 Quiz
  5. Periodicity and Periodic Properties IV | Week 5
    5 Topics
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    1 Quiz
  6. Mass-Volume Relationship in Reaction | Week 6
    8 Topics
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  7. Types of Reactions: Oxidation and Reduction | Week 7 & 8
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  8. Oxidation – Reduction Reaction II | Week 9
    3 Topics
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  9. Electrode Potential and Electrochemical Cells I | Week 10
    6 Topics
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    1 Quiz
  10. Electrode Potential and Electrochemical Cells II | Week 11
    5 Topics
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    1 Quiz
  11. Electrolysis I | Week 12
    8 Topics
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    1 Quiz
  12. Electrolysis II | Week 13
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    1 Quiz

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SS2: CHEMISTRY – 1ST TERM Electrolysis II | Week 13 Theory Questions – Electrolysis II
Lesson 12, Topic 8
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Theory Questions – Electrolysis II

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Topic Content:

  • Theory Questions and Answers – Electrolysis II

Theory Questions – Electrolysis II

1. (a) Sketch to illustrate Faraday’s first law of electrolysis.

(b) Describe how you investigate Faraday’s first law of electrolysis using copper(II) tetraoxosulphate(VI) solution and copper electrodes.

2. Calculate the number of Faraday’s of electricity required to liberate:

(a) 16.2 g of silver 

(b) 14.3 g of aluminum

(c) 2.24 dm3 of oxygen gas at STP

(d) 0.23mole of hydrogen molecule

[Ag = 108, Al = 27, molar volume = 22.4 dm3 at stp]

3. Determine the mass of copper deposited by 4.0 moles of electrons in the reaction represented by the equation: Cu2+ + 2e– → Cu(s) [Cu = 64]

4. How long would it take to deposit 0.08 g of copper from CuCl2 solution by passing a current of 0.5 A?

[Cu = 64, 1F = 96500 C]

5. The following data were recorded in an experiment on electrolysis:-

Table: Data recorded in an experiment on electrolysis 

S/NCurrent (AMPs)Time of current flow (sec)Quantity of electricity (c) Mass of metal (z) deposited (g)
A0.5300______                    0.25
B0.5600______0.65
C0.5_______3750.85
D0.5________4801.30
E0.51500_______1.38
F0.51800_________
 2.28

(a) Copy and complete the table above by calculating the time or quantity of electricity passed in each reaction.

(b) Plot the graph of the mass of Z deposited against the quantity of electricity passed 

(c) From your graph 

(i) How long will it take for the same quantity of electricity to deposit 0.35g of Z?

(ii) Calculate the mass of Z that was deposited by 0.5 A for 27 mins 30 secs 

(iii) What is the quantity of electricity when 1.25 g of the metal is deposited?

(iv) Determine:

  • the electrochemical equivalent 
  • Chemical equivalent 

(v) Which of Faraday’s laws of electrolysis does this experiment Obey?  give the reasons for your answer 

6. In the electrolysis of nickel chloride using carbon electrode, 9.5 g of Nickel was deposited on the cathode electrode. What volume of oxygen will be liberated at STP?

[Ni = 59, if = 96500 C, molar volume = 22.4 d/m3 at stp]

7. Two electrolytic cells containing sodium salt and copper salt electrolytes respectively are connected in series.

A current of 2 Amperes is passed for 30 min. calculate the mass of copper deposited [Na = 23, Cu = 64, 1F = 96500 C]

8. (a) state the following:

(i) Faraday’s first law of electrolysis 
(ii) Faraday’s second law of electrolysis 

(b) State three applications of electrolysis 

(c) A student was asked to electroplate a stainless steel spoon using a brass rod as the electrode and copper (ii) tetraoxosulphate (vi) solution as electrolyte.

The following were the results he obtained 

Mass of spoon = 107.55 g

Mass of copper-coated spoon = ?

Mass of brass rod = 50.4g 

Current = 0.45 A

Time taken = 28 min

(i) Draw a diagram of the setup the student used to electroplate the spoon 

(ii) Which electrode should be the spoon?

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Evaluation Questions

 

1. (a) Sketch to illustrate Faraday’s first law of electrolysis.

Answer:

(b) Describe how you investigate Faraday’s first law of electrolysis using copper(II) tetraoxosulphate(VI) solution and copper electrodes.

Answer:

Different quantities of electricity at various forms are allowed to pass through a fixed concentration of CuSO4 solution using copper electrodes. It could be observed that the higher the quantity of electricity, the greater the amount of copper element deposited at the cathode.

 

2. [Ag = 108, Al = 27, molar volume = 22.4 dm3 at stp]

Calculate the number of Faraday’s of electricity required to liberate:

(a) 16.2 g of silver

Solution

Ag+ + e– → Ag

n = \( \frac{mass}{molar\:mass} \)

= \( \frac{162\:g}{108} \)

= 0.15 moles

1 mole ⇒ 1F

0.15 mole ⇒ \( \frac{0.15 \: \times \: 1}{1} \)

= 0.15 F

 

(b) 14.3 g of aluminum

Solution

Al3+ + 3e– → Al

n = \( \frac{m}{m.m} \\ – \frac{14.3}{27} \\ \scriptsize = 0.5 \)

1 mol ⇒ 3 F

0.5 mol  ⇒ \( \frac{0.5 \: \times \: 3}{1} \)

= 1.5 F

 

(c) 2.24 dm3 of oxygen gas at STP

Solution

O2(g) +4e – → 2O2-

2 mol of  O2 ⇒ 4F

No of moles = \( \frac{volume}{molar\:volume}\)

= \( \frac{2.24}{22.4}\)

= 0.1 mole

2 moles of oxygen ⇒ 4F

0.1 mole ⇒ \( \frac{0.1 \: \times \: 4}{2} \)

= 0.2 F

 

(d) 0.23 mole of hydrogen molecule

Solution

H2(g) → +2e – → 2H+

1 mol ⇒ 2F

0.23 mole ⇒ \( \frac{0.23 \: \times \: 2}{1}\)

= 0.46 F

 

3. Determine the mass of copper deposited by 4.0 moles of electrons in the reaction represented by the equation:

Cu2+ + 2e– →  Cu(s) [Cu = 64]

 

 2F/2 moles of electron ⇒ 1 mol of Cu

 2F/2 moles of electron ⇒ 64 g of Cu

 4 moles of electron ⇒ \( \frac{64 \: \times \: 4}{2} \)

= 128 g

 

4. How long would it take to deposit 0.08 g of copper from CuCl2 solution by passing a current of 0.5 A? [Cu = 64, 1F = 96500 C]

Solution

Cu2+(aq) + 2e–  → Cu(s)

64 g of Cu was deposited by 2F of electricity

64 g of Cu was deposited by (2 × 96500) C of electricity

0.008 g of Cu will deposit \(  \frac{2 \: \times \: 96500 \: \times \: 0.08}{64} \\ = \scriptsize 241.25 C \)

Q = It

t = \(  \frac{Q}{t}\\ =  \frac{241.25}{0.5} \\ = \scriptsize 482.5 \: s \)

convert seconds to minutes

\(  \frac{482.5}{60} \\ = \scriptsize 8.04 \: min \\ \scriptsize \approx 8 \: mins \)

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