Unit 5: Corrosion

Corrosion is the spontaneous destruction or deterioration of a metal by chemical or electrochemical reaction with its environment, converting the refined metal back toward its more stable ore-like state (oxides, hydroxides, sulphides, carbonates). It is essentially the reverse of extractive metallurgy, driven by the tendency of metals to lower their free energy.

Defining features that later sections rely on:

• Thermodynamic driving force: Metals extracted with large energy input are unstable; corrosion releases that energy, so ΔG is negative for a spontaneous corrosion reaction.
• Anodic and cathodic events: Oxidation (loss of electrons, M → M^n+ + ne⁻) always occurs at the anode; reduction consumes those electrons at the cathode.
• Scope and cost: Affects bridges, pipelines, ship hulls, boilers, machinery; causes plant shutdowns, product contamination, loss of strength, and roughly 3–4% of GDP in industrial nations.
• Two broad mechanisms: Direct (dry, chemical) corrosion and electrochemical (wet) corrosion, distinguished by the presence of a conducting liquid.

II. Direct Chemical Corrosion

Dry attack by gases and liquids in the absence of moisture

Chemical corrosion proceeds by direct chemical action of the environment on the metal surface without an electrolyte.

• Oxidation corrosion: Metal reacts with oxygen; 2M + (n/2)O₂ → M₂Oₙ. Governed by the nature of the oxide film formed.
• Pilling–Bedworth rule: Protection depends on the volume ratio of oxide to metal.
  • Stable/protective film: Ratio ≥ 1 and film non-porous, e.g. Al, Cr, Ti form tight oxides that stop further attack.
  • Non-protective film: Ratio < 1; oxide is porous or cracks (e.g. alkali metals like Na), allowing continued oxidation.
• Corrosion by other gases: Cl₂, SO₂, H₂S, CO₂ attack metals; e.g. Cl₂ on Ag gives protective AgCl, whereas on Sn it gives volatile SnCl₄ (accelerated loss).
• Liquid-metal corrosion: Flowing molten metal dissolves solid metal, seen in coolant systems.
• Worked point: Iron above ~500 °C forms layered scale of FeO, Fe₃O₄, Fe₂O₃; the inner porous FeO permits continued oxidation.

III. Electrochemical Corrosion and Different Mechanisms

Wet corrosion through the formation of galvanic cells

Electrochemical corrosion occurs when a metal contacts a conducting electrolyte, setting up separate anodic and cathodic areas connected through the metal and the solution.

• Essential requirements: Anode, cathode, metallic path for electrons, and an electrolyte for ion flow.
• Anodic reaction: Fe → Fe²⁺ + 2e⁻ — metal dissolves, so corrosion always happens at the anode.
• Mechanism 1 — Hydrogen evolution (acidic media): Cathode reduces H⁺; 2H⁺ + 2e⁻ → H₂. Typical of acidic, oxygen-free solutions; whole surface tends to corrode uniformly.
• Mechanism 2 — Oxygen absorption (neutral/alkaline media): Cathode reduces dissolved O₂; O₂ + 2H₂O + 4e⁻ → 4OH⁻. Fe²⁺ and OH⁻ combine, oxidising to rust Fe₂O₃·xH₂O.
• Contrast: Hydrogen-evolution attacks the anodic region continuously; oxygen-absorption localises damage where oxygen access is uneven.

IV. Galvanic and Concentration Cells

Corrosion driven by dissimilar metals or unequal solution conditions

Both types generate an EMF that forces one region to act as anode.

1. Galvanic (bimetallic) corrosion
   • Cause: Two different metals in electrical contact within an electrolyte; the more active metal (higher in the electrochemical/galvanic series) becomes the anode.
   • Example: Steel bolt (cathode) with a zinc washer (anode); zinc corrodes preferentially. In a Zn–Cu couple, Zn dissolves while Cu is protected.
   • Area effect: A small anode with a large cathode gives intense, rapid localised attack.
2. Concentration-cell corrosion
   • Cause: Same metal exposed to electrolyte of differing concentration or aeration; potential difference arises from unequal ion or oxygen levels.
   • Nernst basis: E = E° − (0.0591/n) log[M^n+]; the region contacting dilute metal-ion solution has lower potential and acts as anode.

V. Differential Aeration, Water-line and Pitting Corrosion

Localised attack from uneven oxygen supply

These are special concentration cells where oxygen availability sets up the anode.

A. Differential aeration corrosion

• Principle: The metal area exposed to less oxygen becomes anodic; the well-aerated area becomes cathodic (oxygen supports reduction).
• Consequence: Explains attack under dirt, gaskets, bolt heads, and partly immersed metals.

B. Water-line corrosion

• Site: On tanks and ships storing water, corrosion concentrates just below the waterline.
• Reason: Water surface region is oxygen-rich (cathode); the area a little below has restricted oxygen (anode) and dissolves, forming a distinct corroded band.

C. Pitting corrosion

• Nature: Highly localised attack forming deep, narrow pits, dangerous because small weight loss causes rapid perforation.
• Mechanism: A break in a protective film (scratch, dust particle) creates a small unaerated anode surrounded by a large aerated cathode; Fe²⁺ inside the pit and OH⁻ outside intensify the small-anode/large-cathode effect, deepening the pit autocatalytically.

VI. Intergranular and Soil Corrosion; Factors Affecting Corrosion

Structure-related attack, underground attack, and the variables that control rate

A. Intergranular corrosion

• Definition: Preferential corrosion along grain boundaries of an alloy while grain interiors stay intact.
• Cause: Boundaries are anodic due to segregation or precipitation; in stainless steel, Cr₂₃C₆ precipitates at boundaries, depleting adjacent chromium ("sensitisation") and leaving that zone anodic.
• Effect: Loss of strength and eventual disintegration along grains, common in welds.

B. Soil corrosion

• Definition: Deterioration of buried metal (pipelines, cables) by underground moisture, dissolved salts, and micro-organisms.
• Controlling factors: Soil moisture, porosity (aeration), pH, resistivity, and dissolved chlorides/sulphates; differential aeration between clay and sandy soil sections sets up long-line anodic zones.

C. Factors affecting corrosion

• Nature of the metal:
  • Position in galvanic series: More active metals corrode faster.
  • Purity: Impurities create local cells, so impure metal corrodes faster.
  • Overvoltage and nature of oxide film: High hydrogen overvoltage slows attack; protective (Pilling–Bedworth ≥ 1) films resist.
  • Anode/cathode area ratio: Small anode with large cathode accelerates corrosion.
• Nature of the environment:
  • pH: Acidic media (low pH) increase corrosion; amphoteric metals like Al also corrode in alkali.
  • Temperature: Higher temperature raises reaction and diffusion rates.
  • Humidity and dissolved gases: O₂, CO₂, SO₂ and chlorides raise conductivity and rate.
  • Flow velocity: Faster electrolyte flow removes protective films and supplies more oxygen.

VII. Protection of Corrosion

Reducing attack by design, sacrificial metal, inhibitors, and coatings

Protection works by isolating the metal, altering its potential, or modifying the environment.

A. Protection of corrosion

• Proper design: Avoid dissimilar-metal contact, sharp bends, and crevices that trap moisture; keep anode area large relative to cathode.
• Cathodic protection:
  • Sacrificial anode: A more active metal (Zn, Mg) is attached and corrodes in place of the structure, protecting ship hulls and pipelines.
  • Impressed current: An external DC source forces the structure to become cathodic using an inert auxiliary anode.
• Corrosion inhibitors: Chemicals added to the medium; anodic inhibitors (chromates, phosphates) build protective films, cathodic inhibitors (amines, Mg/Zn salts) suppress the cathodic reaction.
• Metallic coatings:
  • Galvanising (anodic coating): Zn on iron; protects even when scratched because Zn is anodic to Fe.
  • Tinning (cathodic coating): Sn on iron; protects only while intact, since a break makes iron the anode.

B. Protection of corrosion — Ceramic coating

• Principle: Apply a hard, chemically inert non-metallic oxide/silicate layer that isolates the metal from the environment and withstands heat.
• Materials: Metal oxides, silicates, carbides, and glassy enamels (vitreous enamel of borosilicate/alumino-silicate frit fused onto steel).
• Advantages: High-temperature stability, resistance to acids, alkalis and abrasion; used on turbine blades, cookware, and chemical reactors.
• Limitation: Brittle; cracks or chips expose the metal and allow localised corrosion beneath the coat.

VIII. Electroplating Process

Depositing a protective or decorative metal layer by electrolysis

Electroplating coats a conducting object with a thin adherent metal film using an external current to drive reduction of metal ions at the object.

• Cell arrangement:
  • Cathode: The article to be plated (well cleaned/degreased).
  • Anode: A bar of the coating metal (soluble) or an inert anode.
  • Electrolyte: A solution of a salt of the coating metal, often with complexing agents and brighteners.
• Cathode reaction: M^n+ + ne⁻ → M, depositing metal on the article.
• Anode reaction (soluble anode): M → M^n+ + ne⁻, replenishing ions so the bath concentration stays steady.
• Example — copper plating:
  • Electrolyte: CuSO₄ with H₂SO₄.
  • Cathode: Cu²⁺ + 2e⁻ → Cu; Anode: Cu → Cu²⁺ + 2e⁻.
• Controlling variables: Current density, temperature, ion concentration, pH, and additives govern smoothness, adhesion, and grain size of the deposit.
• Purposes: Corrosion resistance, wear resistance, improved appearance, and better solderability; common coatings are Cr, Ni, Cu, Zn, Au, and Ag.