Corrosion Electrochemical Mechanism - UNSOLVED PRACTICE SET
Chapter: Electrochemistry | Topic: Corrosion Electrochemical Mechanism
CORROSION ELECTROCHEMICAL MECHANISM - UNSOLVED PRACTICE SET
Topic: Corrosion Electrochemical Mechanism
Multiple Choice Questions
Q1. The rusting of iron is an example of:
- Dry corrosion
- Wet corrosion (electrochemical corrosion)
- Oxidation in the absence of oxygen
- Physical weathering
Q2. During the rusting of iron, the anodic reaction is:
- Fe → Fe²⁺ + 2e⁻
- O₂ + 2H₂O + 4e⁻ → 4OH⁻
- Fe²⁺ + 2OH⁻ → Fe(OH)₂
- 4Fe(OH)₂ + O₂ → 4FeO(OH) + 2H₂O
Q3. In the electrochemical theory of corrosion, the cathodic reaction involves:
- Loss of electrons (oxidation)
- Gain of electrons (reduction)
- Neutralisation of ions
- Precipitation of hydroxides
Q4. The presence of which of the following accelerates the corrosion of iron?
- Dry air
- Moisture and dissolved salts (like NaCl)
- Pure oxygen without moisture
- Inert gases
Q5. Galvanisation is a method of protecting iron from corrosion by coating it with:
- Copper
- Zinc
- Tin
- Aluminium
Q6. In cathodic protection, the iron structure to be protected is made the:
- Anode
- Cathode
- Electrolyte
- Salt bridge
Short Answer Questions
Q7. Explain the electrochemical mechanism of rusting of iron. Identify the anodic and cathodic reactions that occur in the presence of water and oxygen.
Q8. Why does iron rust faster in coastal areas (like Chennai or Mumbai) compared to dry inland areas (like Jaipur)? Explain in terms of the electrochemical theory of corrosion.
Q9. What is galvanisation? Explain how coating iron with zinc protects it from corrosion, even if the zinc layer is scratched.
Q10. Differentiate between 'dry corrosion' and 'wet corrosion' (electrochemical corrosion). Give one example of each.
Q11. What is cathodic protection? Explain sacrificial anode protection with a suitable example. Why is magnesium or zinc used as the sacrificial anode?
Q12. Write the overall chemical reaction for the formation of rust (hydrated ferric oxide, Fe₂O₃·xH₂O) from iron. How many moles of oxygen are required to form 1 mole of rust from Fe?
Long Answer Questions
Q13. (a) Explain the electrochemical theory of corrosion with reference to the rusting of iron. Draw a neat diagram showing the formation of an electrochemical cell on the surface of iron.
(b) Identify the anodic and cathodic areas, and write the half-reactions occurring at each.
(c) How does the presence of dissolved CO₂ in water affect the rate of corrosion?
Q14. (a) What is meant by 'cathodic protection'? Describe the two methods of cathodic protection: (i) sacrificial anode method, and (ii) impressed current method.
(b) Why is the sacrificial anode method particularly useful for protecting underground pipelines and ship hulls?
(c) Name two metals commonly used as sacrificial anodes and explain why they are chosen.
Q15. (a) Explain the following methods of preventing corrosion: (i) Barrier protection (painting, oiling, greasing), (ii) Galvanisation, and (iii) Alloying.
(b) Why is stainless steel (an alloy of iron with chromium and nickel) more resistant to corrosion than pure iron?
(c) Why do iron pipes carrying hot water corrode faster than those carrying cold water?
Numerical / Application-Based Problems
Q16. An iron pipeline in a coastal city in India (like Chennai or Kochi) is found to be corroding rapidly. The pipeline is exposed to moist air and salt water spray from the sea.
(a) Explain the electrochemical mechanism of corrosion occurring on this iron pipeline. Write the half-reactions for the anodic and cathodic processes, and the overall reaction leading to the formation of rust (Fe₂O₃·xH₂O).
(b) Suggest and explain two practical methods that can be used to protect this pipeline from further corrosion. Your answer should include the scientific principle behind each method.
(c) Why does corrosion occur more rapidly in coastal areas compared to dry inland regions? Explain in terms of conductivity and oxygen availability.
Q17. A steel bridge support in Mumbai is protected using sacrificial anodes made of magnesium. The support has a surface area of 500 m² and requires protection for 10 years.
(a) Explain why magnesium is chosen as the sacrificial anode rather than copper or tin.
(b) If the corrosion rate of iron in the absence of protection is 0.5 mm per year, calculate the total thickness of iron that would be lost in 10 years without protection.
(c) The standard electrode potentials are: E°(Mg²⁺/Mg) = −2.37 V, E°(Fe²⁺/Fe) = −0.44 V, E°(Cu²⁺/Cu) = +0.34 V. Using these values, explain why magnesium protects iron but copper would accelerate its corrosion.
(d) Write the reactions occurring at the magnesium anode and the iron cathode in this protection system.
Q18. A student observes that an iron nail placed in a test tube with boiled, deaerated water and a layer of oil on top does NOT rust, while another iron nail placed in a test tube with ordinary tap water rusts within a few days.
(a) Explain why the nail in boiled, deaerated water with oil does not rust. What role does each of the following play: (i) boiling the water, (ii) the oil layer?
(b) Design an experiment to prove that BOTH oxygen and water are necessary for the rusting of iron. Describe the setup, the observations expected, and the conclusion drawn.
(c) If 5.6 g of iron completely corrodes to form Fe₂O₃·xH₂O (rust), calculate the maximum mass of rust that can be formed. Assume x = 3 for simplicity.
(Given: Atomic masses: Fe = 56, O = 16, H = 1 u)