Corrosion Notes

Unit VI: Corrosion

Basic Concepts

• Basic Corrosion Cell: Consists of an anode, cathode, metallic path, and electrolyte.
• Anode: Where oxidation occurs (e.g., Fe → Fe²⁺ + 2e⁻).
• Cathode: Where reduction occurs (e.g., 2H⁺ + 2e⁻ → H₂ or O₂ + 2H₂O + 4e⁻ → 4OH⁻).
• Metallic Path: Provides a route for electron flow between the anode and cathode.
• Electrolyte: A conducting medium that facilitates ion transport (e.g., water containing ions).
• Rust: A corrosion product, typically iron oxide (Fe₂O₃·nH₂O).

Definition of Corrosion

• Corrosion is the deterioration of materials through chemical interaction with their environment.
• Primarily refers to metals but can also apply to plastics, concrete, and wood.

Detailed Corrosion Definition

• Corrosion: Unintentional destruction of a solid material by chemical or electrochemical reactions starting from the surface.
• Non-metals corrode via chemical reactions.
• Metals corrode mostly via electrochemical reactions.
• Metals tend to revert to a lower energy state by forming oxides, hydroxides, carbonates, or sulphides.

Causes of Corrosion

• Metals are extracted from ores, which are in a chemically combined state.
• Ore: The chemical combined state of the metal with low energy.
• Extracted metals possess high energy and are thermodynamically unstable.
• Metals corrode to achieve a stable state by reacting with the environment.
• Corroded metal is thermodynamically stable.

Rusting of Iron in Acidic Environment (Oxygen Absent)

• Large Anodic Area: $Fe → Fe^{2+} + 2e^-$
• Small Cathode: $2H^+ + 2e^- → H_2$
• Overall Reaction: $Fe + 2H^+ → Fe^{2+} + H_2$

Effects/Disadvantages of Corrosion

• Loss of valuable metallic properties (conductivity, malleability, ductility).
• Reduced lifespan of metallic machinery parts.
• Enormous wastage of metal in compound form.
• Machinery failure due to loss of metal properties.
• Significant economic loss: approximately $2 to 2.5 billion per annum worldwide.

Theories/Mechanisms of Corrosion

• Dry or Chemical Corrosion
• Wet or Electrochemical Corrosion

Dry Corrosion (Chemical Corrosion)

• Direct chemical reaction of environment gases or inorganic liquids with metal surfaces.
• Occurs due to direct chemical action of the environment on the metal surface without moisture or liquid electrolyte.
• Typically occurs at temperatures above 100°C.
• Three types:
  • Oxidation Corrosion
  • Corrosion by Other Gases
  • Liquid Metal Corrosion

Factors Affecting Dry Corrosion

1. Chemical affinity between the environment and the metal.
2. Protective value of the formed film.
3. Nature of the film formed.
4. Adhesion between the film and the metal surface.

Oxidation Corrosion

• Direct action of oxygen at low or high temperatures in the absence of moisture.
• Corrosion rate depends on temperature.
• Alkali and alkaline earth metals oxidize rapidly at low temperatures.
• At high temperatures, all metals except Ag, Au, and Pt oxidize.
• Oxidation: $M → M^{2+} + 2e^-$
• Reduction: $O_2 + 2e^- → 2O^-$
• Metal Oxide Formation: $M + O_2 → M^{2+} + 2O^-$

Oxidation Corrosion Mechanism

1. Metal oxidation (loss of electrons): $M → M^{+n} + ne^-$
2. Oxygen gains electrons to form oxide ions: $nO_2 + ne^- → 2nO^-$
3. Metal oxide scale formation: $M + nO_2 → M^{n+} + 2nO^-$

Oxide Layer Cases

1. Stable Oxide Layer: Fine-grained, tightly adhering, and impervious (e.g., Al, Sn, Pb, Cu). Acts as a protective layer.
2. Unstable Oxide Layer: Oxide decomposes back into metal and oxygen (e.g., Ag, Pt, Au). No corrosion.
3. Volatile Oxide Layer: Oxide volatilizes as it forms, exposing fresh metal (e.g., molybdenum). Rapid corrosion.
4. Porous Oxide Layer: Oxide film has pores/cracks, allowing O₂ to penetrate (e.g., iron and steel). Continuous corrosion.

Corrosion Due to Other Gases

• Gases like SO₂, CO₂, Cl₂, H₂S, F₂ cause corrosion.

• Corrosive effect depends on the chemical affinity between the metal and the gas.

• Attack depends on the formation of protective or non-protective films.

  • Protective film: If the volume of the corrosion film is greater than the underlying metal, it is strongly adherent and non-porous, preventing further corrosive gas penetration. Example: $Ag + Cl_2 → 2AgCl$
  • Non-protective film: If the volume of the corrosion film is less than the underlying metal or it forms pores/cracks, allowing corrosive gas penetration. Example:
    • $SnCl_4$ is volatile, causing excessive corrosion of Sn.
    • $H_2S$ attacks steel, forming porous $FeS$ layer.

Liquid Metal Corrosion

• Anhydrous liquid attacks the metal surface.
• Liquid metal flowing over solid metal at high temperature weakens it due to:
  • Dissolution in liquid metal.
  • Penetration of liquid metal into solid metal.
• Example: Sodium (coolant) corrodes cadmium in a nuclear reactor.

Wet or Electrochemical Corrosion

• Occurs when a conducting liquid is in contact with a metal or when dissimilar metals are immersed in a solution.
• Corrosion occurs due to anodic and cathodic areas.
• Anode: Oxidation reactions.
• Cathode: Reduction reactions.
• Requires aqueous solution or liquid electrolytes.
• Most efficient in waters containing salts (e.g., NaCl in marine conditions).

Electrochemical Corrosion Mechanism

• Anodic Reaction: Metal dissolution: $M ↔ M^{+n} + ne^-$
• Cathodic Reactions:
  • Hydrogen gas evolution
  • Oxygen gas absorption

Wet Corrosion by Hydrogen Evolution

• Occurs when anodic areas are large, cathodic areas are small, and oxygen is absent.
• Acidic Medium:
  • Anode: $Fe → Fe^{2+} + 2e^-$
  • Cathode: $2H^+ + 2e^- → H_2$
  • Overall: $Fe + 2H^+ → Fe^{2+} + H_2$
• All metals above hydrogen in the electrochemical series dissolve in acidic solution, evolving H₂ gas.

Wet Corrosion in Neutral or Alkaline Medium

• Anodic reaction: $Fe → Fe^{2+} + 2e^-$
• Cathodic reaction: $2H<em>2O + 2e^- → H</em>2 + 2OH^-$
• Overall reaction: $Fe + 2H<em>2O → Fe^{2+} + H</em>2 + 2OH^-$
• $Fe^{2+}$ and $OH^-$ meet to form $Fe(OH)_2$.

Rusting of Iron in Neutral Aqueous Solution

• Iron surface coated with a thin film of iron oxide.
• Cracks in the film create anodic areas.
• Metal parts act as cathodes.
• Anodes are small; cathodes are large.
• Electrons flow from anode to cathode through iron.

Absorption of Oxygen

• Anodic areas are small, and cathodic areas are large.
• Oxygen is present, and the environment is neutral or alkaline.
• Corrosion product forms closer to the cathode.

Rust Formation

• Anode: $Fe → Fe^{2+} + 2e^-$
• Cathode: $\frac{1}{2}O<em>2 + H</em>2O + 2e^- → 2OH^-$
• Overall: $Fe^{2+} + 2OH^- → Fe(OH)_2$
• $4Fe(OH)<em>2 + O</em>2 + 2H<em>2O → 4Fe(OH)</em>3$ or $2Fe<em>2O</em>3 · 3H_2O$ (Rust)

Chemical vs. Electrochemical Corrosion

Types of Electrochemical Corrosion

• Differential Metallic Corrosion (DMC) or Bimetallic corrosion
  • Galvanic corrosion
• Concentration cell corrosion
  • Pitting corrosion
  • Waterline corrosion
  • Stress corrosion
  • Corrosion under a drop of water
  • Caustic embrittlement

Bimetallic (Galvanic) Corrosion

• Occurs when different metals are in contact and exposed to a corrosive atmosphere.
• The metal with a higher electrode potential value forms the anode and corrodes.
• Example: In a Zn-Cu galvanic cell, zinc is the anode and corrodes; copper is the cathode and is protected.
• Acidic environment: $2H^+ + 2e^- → H_2$
• Alkaline/neutral environment: $\frac{1}{2}O<em>2 + H</em>2O + 2e^- → 2OH^-$
• $Zn^{2+}$ and $2OH^-$ form $Zn(OH)_2$.
• Examples:
  • Cu pipes with iron pipes
  • Steel propeller shaft in a bronze bearing
  • Steel screw in brass marine hardware
  • Lead-antimony solder around Cu wire

Electrochemical Series

• Lists metals in order of their reducing activity (tendency to lose electrons).
• Metals higher in the series are more likely to corrode.
• Example pairs and their equilibrium potentials (E'):
  • Lithium: $Li^+ + e^- → Li(s)$, E' = -3.03 V
  • Potassium: $K^+ + e^- → K(s)$, E' = -2.92 V
  • Calcium: $Ca^{2+} + 2e^- → Ca(s)$, E' = -2.87 V
  • Sodium: $Na^+ + e^- → Na(s)$, E' = -2.71 V
  • Magnesium: $Mg^{2+} + 2e^- → Mg(s)$, E' = -2.37 V
  • Aluminum: $Al^{3+} + 3e^- → Al(s)$, E' = -1.66 V
  • Zinc: $Zn^{2+} + 2e^- → Zn(s)$, E' = -0.76 V
  • Iron: $Fe^{2+} + 2e^- → Fe(s)$, E' = -0.44 V
  • Lead: $Pb^{2+} + 2e^- → Pb(s)$, E' = -0.13 V
  • Hydrogen: $2H^+ + 2e^- → H_2(g)$, E' = 0.00 V
  • Copper: $Cu^{2+} + 2e^- → Cu(s)$, E' = +0.34 V
  • Silver: $Ag^+ + e^- → Ag(s)$, E' = +0.80 V
  • Gold: $Au^{3+} + 3e^- → Au(s)$, E' = +1.50 V

Concentration Cell (Differential Aeration) Corrosion

• Occurs when one part of the metal is exposed to a different air concentration than another.
• Potential difference between differently aerated areas.
• Poorly oxygenated areas become anodic.
• Highly oxygenated areas become cathodic.

Pitting Corrosion

• Explained by differential aeration.
• The pit becomes deeper, and its bottom becomes less open to oxygen, making it more anodic.
• Localized corrosion in pits, cavities, and pinholes.
• Oxygen-deficient pit acts as the anode; the plane surface is the cathode.
• Small anode area results in a high corrosion rate.
• Pitting is common in aluminum alloys, copper alloys, stainless steels, and some nickel alloys.
• Pits are initiated by activating ions like chloride ions.

Waterline Corrosion

• Oxygen concentration is higher at the water surface than deeper down.
• Forms a concentration cell.
• The lower portion is the anode.
• The surface at the water level is the cathode.
• If Zn is partially immersed in NaCl solution:
  • areas above, close to the waterline are well aerated and act as cathode
  • areas deep inside the solution are anodic, as [O₂] is less; potential develops, causing current flow between the two areas of the same metal; Zn dissolves at the anodic area

Intergranular Corrosion

• Occurs along grain boundaries.
• Grain boundaries act as anodes due to precipitation of certain compounds.
• Precipitated compounds and grain centers behave as cathodes.

Passivity

• The phenomenon where a metal exhibits higher corrosion resistance than expected.
• Some metals, like Ti and Al, develop strongly adhering oxide layers, making their effective electrode potential more positive (less negative).
• Passivity is due to a thin (0.004-mm-thick) invisible oxide film.

Factors Affecting Corrosion

• Nature of the metal
• Nature of the environment

Nature of the Metal

• Purity of the Metal: Impurities create electrochemical cells, leading to corrosion at anodic parts.
• Electrode Potentials: Metals with higher reduction potentials (noble metals) corrode less easily.
• Position in Galvanic Series: Metals higher in the series corrode more easily.
• Relative Areas of Anodic and Cathodic Cells: Corrosion rate is proportional to the cathodic/anodic area ratio.
• Physical State of Metal: Metals with small grain size corrode more.
• Hydrogen Overvoltage: Metals with lower hydrogen overvoltage are more susceptible to corrosion.
• Nature of Surface Film: Stable, insoluble, and non-porous films protect against corrosion.

Nature of the Environment

• Temperature: Corrosion rate increases with temperature.
• Humidity in the Air: Moisture provides water for the electrolyte.
• Presence of Impurities: Gases like CO₂, SO₂, H₂S increase corrosion.
• pH Value: Acidic pH increases the rate of corrosion.
• Amount of Oxygen in Atmosphere: Higher oxygen levels increase corrosion.
• Velocity of Ions: Higher diffusion rates increase corrosion.

Corrosion Control Methods

• Cathodic Protection
• Surface Coatings

Cathodic Protection

• Protecting metals by making them cathodes.
• An auxiliary anode is provided in the corroding medium and connected to the structure.
• Types:
  • Sacrificial Anodic Method: Connecting the metal to a more anodic metal (e.g., Mg, Zn, Al).
  • Impressed Current Cathodic Method: Applying direct current to nullify the corrosion current.

Sacrificial Anodic Protection

• Metal to be protected is connected to a more anodic metal.
• Commonly used metals: Mg, Zn, Al, and their alloys.
• Example: A ship hull (steel) connected to Zn blocks, which corrode, protecting the steel.
• Example: Underground water pipelines and water tanks are protected with sacrificial anodes.

Impressed Current Method

• Direct current applied in the opposite direction to nullify the corrosion current.
• Converts the corroding metal from anode to cathode.