Corrosion and its control

Introduction to Corrosion

Corrosion is the degradation of metals due to reactions with surrounding environments (gases or liquids).
Common examples of corrosion:
- Rusting of Iron: Formation of reddish-brown scale (Fe2O3.x H2O).
- Copper Tarnishing: Green layer of basic copper carbonate [CuCO3+Cu(OH)2].
- Silver Tarnishing: Black layer of silver sulphide formed.

Most metals, except gold and platinum, corrode spontaneously when exposed to air.
Metals exist as combined states in ores (oxides, carbonates, sulphides) which are more stable.
High energy is required to extract pure metals from ores, making them thermodynamically unstable.
Metals naturally revert to a stable state when in contact with elements in the environment.

Ore of Metal (stable state) → Metal extraction + Energy → Pure Metal (unstable state) → Corrosion → Corroded Metal + Energy (more stable than pure).

Corrosion can be classified into:
- Chemical or Dry Corrosion.
- Electrochemical or Wet Corrosion.

Occurs in the presence of dry gases (O2, N2, H2S, etc.) or anhydrous liquids due to direct chemical reactions.
Types include:
- Oxygen Corrosion: Involves direct attack by oxygen.
- Corrosion by Other Gases: Affected by gases like Cl2 and H2S.
- Liquid Metal Corrosion: Occurs at high temperatures with metal flowing over another.

Happens in moist conditions or electrolyte solutions.
Key conditions for electrochemical corrosion include:
- Presence of cathodic and anodic areas.
- Electrical potential across areas.
- A conductive metallic path is essential.

Metal in moist environments forms electrochemical cells on its surface:
- At the anode: Metal oxidizes and produces metal ions (e.g., Fe → Fe2+ + 2e−).
- At the cathode: Reduction reactions occur leading to rust formation.
Example: Rusting of iron involves:
- Anodic reaction: Iron oxidizes to ferrous ions.
- Cathodic reactions can vary based on conditions (e.g., presence of oxygen or hydrogen).

Hydrogen Evolution: Occurs in acidic environments.
Absorption of Oxygen: Occurs in neutral/alkaline conditions, resulting in the formation of rust.

Chemical Corrosion: Occurs under dry conditions; direct chemical reaction.
Electrochemical Corrosion: Takes place in wet conditions; involves galvanic cell formation.

Metals exhibit passivity when a stable oxide layer builds on their surface, acting as a barrier to further corrosion.
Metals like zirconium and aluminum form protective oxide films in ambient conditions.

The series predict corrosion tendencies; metals are arranged from most to least noble (e.g., Mg to Pt).

Galvanic Corrosion: Occurs when dissimilar metals are in contact in a conductive environment.
Differential Aeration Corrosion: Variations in oxygen levels lead to localized corrosion.
Pitting Corrosion: Localized attack causes deep holes in metals.
Stress Corrosion: Caused by tensile stress in conjunction with corrosive environments.

Nature of Metal: Position in galvanic series, purity, anodic/cathodic area ratio, and the nature of oxide films.
Environmental Factors: Temperature, humidity, and impurity presence enhance corrosion rate.

Proper Designing of Materials: Avoid conductive connections between dissimilar metals, ensure uniform aeration and minimize crevices.
Material Selection: Use of pure metals or resistant alloys.
Cathodic Protection:
- Sacrificial Anode Protection: Connection to a more reactive metal to protect the target metal.
- Impressed Current Protection: External current applied to suppress corrosion.
Anodizing: Electrochemical process to improve the protective layer on metals like aluminum.
Corrosion Inhibitors: Categories include anodic (films on anodic surfaces) and cathodic inhibitors (adsorption on cathodic sites).

Metallic Coatings: e.g., galvanization (zinc coating) prevents rusting.
Non-metallic Coatings: Physical barriers; can be paint or other protective layers.