Zinc is used to protect the iron from corrosion because zinc is
more electropositive than iron
Protecting Iron from Corrosion: Why Zinc Works
Iron is a widely used metal, but it is susceptible to rusting, which is a form of corrosion. Corrosion is the degradation of a material due to reactions with its environment. Rusting specifically refers to the oxidation of iron in the presence of water and oxygen, forming iron oxides (rust).
To protect iron structures from this damaging process, various methods are employed. One very common and effective method is coating the iron with a layer of zinc. This process is called galvanization.
Understanding Galvanization and Sacrificial Protection
Galvanization involves dipping the iron object into molten zinc or electroplating it with zinc. The resulting zinc coating acts as a barrier, preventing oxygen and water from reaching the iron surface. However, the protection offered by zinc is more than just a physical barrier. Even if the zinc layer is scratched, exposing the iron underneath, the iron still does not rust immediately.
This enhanced protection is due to a concept called sacrificial protection. Zinc is used because of its electrochemical properties relative to iron.
Comparing Electrochemical Properties of Zinc and Iron
Metals have different tendencies to lose electrons and become positive ions. This tendency is related to their electropositivity or their position in the electrochemical series.
- A more electropositive metal has a greater tendency to lose electrons and corrode (oxidize).
- A less electropositive metal has a lower tendency to lose electrons.
When zinc and iron are in contact in the presence of an electrolyte (like moisture), they form an electrochemical cell. In this cell, the more electropositive metal acts as the anode and gets oxidized (corrodes), while the less electropositive metal acts as the cathode and is protected.
| Metal | Tendency to Lose Electrons | Behavior When Coupled with a Less Electropositive Metal |
|---|---|---|
| Zinc (Zn) | Higher (More Electropositive) | Acts as anode, corrodes |
| Iron (Fe) | Lower (Less Electropositive) | Acts as cathode, protected |
Because zinc is more electropositive than iron, in the presence of moisture, zinc corrodes preferentially, sacrificing itself to protect the iron. The zinc atoms lose electrons and turn into zinc ions, while the iron remains in its metallic state, thus preventing rust formation.
Analyzing the Options
Let's look at the given options in the context of why zinc protects iron from corrosion:
- more electropositive than iron: This property is the key reason for zinc's ability to provide sacrificial protection to iron. Because zinc is more electropositive, it corrodes preferentially, protecting the iron. This aligns with the principle of galvanization and sacrificial protection.
- cheaper than iron: While cost can be a factor in choosing materials for large-scale applications, it is not the fundamental chemical or physical reason why zinc provides corrosion protection to iron. Zinc is not necessarily cheaper than iron; its use is based on its protective properties.
- a bluish white metal: This is a physical characteristic of zinc (its colour and appearance). While true, the colour or appearance of zinc has no bearing on its ability to prevent iron from rusting through an electrochemical process.
- a good conductor of heat and electricity: Zinc is indeed a good conductor, but this property is not the reason it protects iron from corrosion. Corrosion is an electrochemical process, but zinc's conductivity isn't the protective mechanism; its relative electropositivity is.
Therefore, the primary reason zinc is used to protect iron from corrosion is its electrochemical nature – being more electropositive than iron, allowing it to act as a sacrificial anode.
Revision Table: Key Concepts
| Term | Definition/Significance |
|---|---|
| Corrosion | Degradation of a material due to reaction with its environment (e.g., rusting of iron). |
| Rusting | Specific term for the corrosion of iron (oxidation of iron). |
| Galvanization | Coating iron with a layer of zinc to protect it from corrosion. |
| Electropositivity | Tendency of an atom to lose electrons and form positive ions. More electropositive metals are more reactive and corrode easily. |
| Sacrificial Protection | A method of corrosion prevention where a more electropositive metal (sacrificial anode) is deliberately connected to the metal to be protected. The sacrificial metal corrodes preferentially. |
Additional Information: Applications and Alternatives
Galvanized iron is widely used in roofing sheets, pipes, fences, and other outdoor structures where protection from moisture and air is essential. The sacrificial action of zinc means that even if the coating is damaged, the iron beneath is still protected as long as some zinc is present and in electrical contact.
Other methods to protect iron from corrosion include painting, oiling, greasing, tin plating, chrome plating, and alloying (like making stainless steel by adding chromium and nickel). However, galvanization is particularly effective due to the active sacrificial protection offered by zinc.
Important Questions from Electrochemistry
Which one of the following is the correct arrangement of metals in the decreasing order of their reactivity?
Which one of the following represents the correct order of electron releasing tendency of metals?
Which compound, when dissolved in water, conducts electricity and forms a basic solution?
Identify an application in which solar cells are NOT used.
In the symbol of the electric cell, the thicker, shorter line represents the______.
