Corrosion (Electrochemical Theory)

Corrosion

Corrosion: Corrosion is the gradual destruction of materials (usually metals) by chemical or electrochemical reaction with their environment. It is a natural process that leads to the degradation of materials over time.

Economic Impact: Corrosion causes significant economic losses worldwide due to the cost of replacing corroded materials, maintaining structures, and implementing protective measures.

Electrochemical Theory of Corrosion: The most widely accepted explanation for the corrosion of metals is based on electrochemical principles, similar to those occurring in a galvanic cell.

Process of Rusting of Iron (A Common Example):

Rusting is the corrosion of iron and its alloys, forming iron oxides.

Electrochemical Cell Formation: Corrosion typically involves the formation of tiny electrochemical cells on the surface of the metal.

1. Anodic Areas: Certain areas on the metal surface act as anodes where oxidation (loss of electrons) occurs. For iron, this is typically where the metal is less pure or stressed.
2. Cathodic Areas: Other areas act as cathodes where reduction occurs.
3. Electrolyte: A thin film of moisture on the metal surface acts as an electrolyte, allowing ion movement. This moisture often contains dissolved substances like acids or salts that increase its conductivity.

Steps in Rusting of Iron:

1. Anodic Reaction (Oxidation): Iron metal loses electrons and gets oxidized to ferrous ions ($Fe^{2+}$).

$$Fe(s) \rightarrow Fe^{2+}(aq) + 2e^- \quad (E^\circ = -0.44 \text{ V})$$

2. Cathodic Reaction (Reduction): Electrons released at the anode travel through the metal to the cathodic areas. At the cathode, oxygen dissolved in the moisture acts as the oxidizing agent and gets reduced.

In neutral or alkaline solutions: $O_2(g) + 2H_2O(l) + 4e^- \rightarrow 4OH^-(aq) \quad (E^\circ = +0.40 \text{ V})$

In acidic solutions: $O_2(g) + 4H^+(aq) + 4e^- \rightarrow 2H_2O(l) \quad (E^\circ = +1.23 \text{ V})$

3. Formation of Rust: The ferrous ions ($Fe^{2+}$) formed at the anode migrate through the electrolyte layer and are further oxidized by dissolved oxygen to ferric ions ($Fe^{3+}$).

$$4Fe^{2+}(aq) + O_2(g) + 4H^+(aq) \rightarrow 4Fe^{3+}(aq) + 2H_2O(l)$$

4. The ferric ions ($Fe^{3+}$) then react with water to form hydrated ferric oxide, which is rust. Rust is typically represented as $Fe_2O_3 \cdot nH_2O$.

$$2Fe^{3+}(aq) + 6H_2O(l) \rightarrow 2Fe(OH)_3(s) \rightarrow Fe_2O_3 \cdot nH_2O(s) + (3-n)H_2O(l)$$

Factors Accelerating Corrosion:

• Presence of Moisture: Essential for forming the electrolyte layer.
• Presence of Dissolved Oxygen: Acts as the cathodic reactant.
• Presence of Electrolytes (Salts, Acids): Dissolved salts (like $NaCl$) and acids increase the conductivity of the moisture film, accelerating the electrochemical process. For example, rusting is faster in saltwater than in pure water.
• Impurities: Impurities in metals can create galvanic cells, leading to localized corrosion. For example, contact between iron and less reactive metals like copper or tin can accelerate iron's corrosion.
• Stress and Strain: Stressed areas of a metal can become anodic and corrode more rapidly.

Prevention of Corrosion: Methods to prevent corrosion are aimed at interrupting the electrochemical process:

• Protective Coatings: Applying a barrier layer to prevent contact with the corrosive environment.
• Painting or Varnishing: Common for preventing rust.
• Oiling or Greasing: Forms a protective film.
• Galvanization: Coating iron or steel with a layer of zinc. Zinc is more reactive than iron, so it corrodes preferentially (acts as a sacrificial anode), protecting the iron. Even if the zinc coating is scratched, it continues to protect the iron electrochemically.
• Chromium Plating: Provides a protective and decorative layer.
• Sacrificial Protection: Attaching a more reactive metal (sacrificial anode) to the metal to be protected. The sacrificial anode corrodes instead of the protected metal.
• Example: Attaching blocks of magnesium or zinc to steel pipelines or ship hulls.
• Alloying: Forming alloys where the metal is mixed with other elements to enhance its resistance to corrosion.
• Stainless Steel: An alloy of iron, chromium, nickel, and carbon, which forms a passive protective layer of chromium oxide on its surface.
• Anodic Protection: Passivating the metal surface by forming a protective oxide layer, often by making the metal the anode in an electrochemical cell.