Corrosion: Mechanism and Prevention

Introduction

Every year, corrosion costs the global economy over $2.5 trillion - about 3.4% of global GDP! From the Titanic rusting at the ocean floor to the Golden Gate Bridge requiring constant maintenance, corrosion is everywhere. Understanding this spontaneous electrochemical process is crucial not just for JEE, but for protecting the infrastructure of modern civilization!

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Interactive: Corrosion Process Animation

Watch how iron corrodes through electrochemical reactions:

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What is Corrosion?

Corrosion is the gradual destruction of a metal by spontaneous electrochemical reaction with its environment, converting refined metal back to its natural oxide/salt form.

Why Corrosion Occurs

Metals in nature exist as ores (oxides, sulfides, carbonates). Extracting pure metal requires energy input:

$$\text{Fe}_2\text{O}_3 \xrightarrow{\text{+Energy}} \text{Fe} + \text{O}_2$$

Pure metal has higher energy than the ore. Nature wants to return to the lower energy state:

$$\text{Fe} + \text{O}_2 \xrightarrow{\text{Spontaneous}} \text{Fe}_2\text{O}_3$$

This spontaneous process is corrosion!

Corrosion as an Electrochemical Process

Corrosion involves:

1. Oxidation of metal (metal loses electrons)
2. Reduction of environmental species (O₂, H⁺ gain electrons)
3. Flow of electrons through the metal
4. Formation of corrosion products (oxides, hydroxides)

It’s essentially a galvanic cell operating on the metal surface!

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Rusting of Iron

Rusting is the corrosion of iron and steel, forming hydrated ferric oxide (Fe₂O₃·xH₂O), commonly called rust.

Conditions Required for Rusting

1. Oxygen (O₂) from air
2. Water (H₂O) or moisture
3. Electrolyte (salts, acids in water increase rate)

Note: Pure, dry air alone won’t cause rusting. Pure water alone is very slow. Both are needed!

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Mechanism of Rusting

Rusting occurs through formation of tiny electrochemical cells on the iron surface.

Step-by-Step Process

Step 1: Formation of Anodic and Cathodic Regions

On iron surface, different regions become anode and cathode due to:

• Surface impurities
• Grain boundaries
• Stress differences
• Concentration variations

At Anode (oxidation):

$$\boxed{\text{Fe}_{(s)} \rightarrow \text{Fe}^{2+}_{(aq)} + 2e^-}$$

Iron atoms lose electrons and go into solution as Fe²⁺ ions. This creates pits in the metal.

At Cathode (reduction):

In acidic/neutral conditions:

$$\text{O}_2 + 4\text{H}^+ + 4e^- \rightarrow 2\text{H}_2\text{O}$$

In basic/neutral conditions (more common):

$$\boxed{\text{O}_2 + 2\text{H}_2\text{O} + 4e^- \rightarrow 4\text{OH}^-}$$

Oxygen dissolved in water layer accepts electrons.

Step 2: Electron Flow

Electrons flow from anodic to cathodic regions through the metal itself!

Step 3: Formation of Fe(OH)₂

Fe²⁺ ions migrate toward cathodic region and meet OH⁻ ions:

$$\text{Fe}^{2+} + 2\text{OH}^- \rightarrow \text{Fe(OH)}_2$$

Ferrous hydroxide forms as a greenish precipitate.

Step 4: Oxidation to Fe(OH)₃

Fe(OH)₂ is further oxidized by oxygen:

$$4\text{Fe(OH)}_2 + \text{O}_2 + 2\text{H}_2\text{O} \rightarrow 4\text{Fe(OH)}_3$$

Step 5: Dehydration to Rust

$$2\text{Fe(OH)}_3 \rightarrow \text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O} + (3-x)\text{H}_2\text{O}$$

This reddish-brown hydrated ferric oxide is rust!

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Overall Rusting Equations

Complete reaction (simplified):

$$\boxed{4\text{Fe} + 3\text{O}_2 + 2x\text{H}_2\text{O} \rightarrow 2\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}}$$

In presence of acids/CO₂:

$$2\text{Fe} + \text{O}_2 + 4\text{H}^+ \rightarrow 2\text{Fe}^{2+} + 2\text{H}_2\text{O}$$

followed by oxidation to Fe³⁺ and precipitation.

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Factors Affecting Rate of Corrosion

1. Nature of Metal

Position in electrochemical series:

• Lower E° (more negative) = more reactive = faster corrosion
• Higher E° = less reactive = slower corrosion

Reactivity order for corrosion:

$$\text{Na, K, Mg} > \text{Al} > \text{Zn} > \text{Fe} > \text{Ni} > \text{Sn} > \text{Pb} > \text{H} > \text{Cu} > \text{Ag} > \text{Au}$$

Examples:

• Sodium corrodes instantly in air (stored in kerosene)
• Aluminum forms protective oxide layer (passivation)
• Gold doesn’t corrode (King Tut’s mask after 3300 years!)

2. Presence of Impurities

Pure metals corrode slower than alloys or impure metals.

Why?

• Impurities create galvanic cells on surface
• Different metals have different electrode potentials
• Electrons flow from more reactive to less reactive metal

Example: Iron with carbon impurities

• Fe (E° = -0.44 V) acts as anode → corrodes
• C (inert) acts as cathode → protected
• Iron corrodes faster near carbon!

3. Presence of Electrolyte

More ions in water = faster corrosion

Examples:

• Seawater: High salt content → very corrosive (ship hulls!)
• Rainwater with CO₂/SO₂: Acidic → accelerates corrosion
• Pure distilled water: Slow corrosion (low conductivity)

Acid rain effect:

$$\text{SO}_2 + \text{H}_2\text{O} \rightarrow \text{H}_2\text{SO}_3$$ $$2\text{NO}_2 + \text{H}_2\text{O} \rightarrow \text{HNO}_3 + \text{HNO}_2$$

These acids dramatically increase H⁺ concentration, accelerating cathode reaction!

4. Temperature

Higher temperature = faster corrosion (generally)

Reasons:

• Increased reaction rates (Arrhenius equation)
• More water evaporation → concentration increases
• But: Very high temperatures can reduce dissolved O₂

5. pH of Environment

pH Range	Corrosion Rate	Example
pH < 4	Very high	Acid rain, industrial areas
pH 4-10	Moderate	Normal atmospheric conditions
pH > 10	Low	Alkaline environments (protective films)

Acidic conditions:

• More H⁺ available for cathode reaction
• Dissolves protective oxide layers
• Iron corrodes rapidly

Alkaline conditions:

• Forms protective hydroxide films
• Slows corrosion
• Used in some protective treatments

6. Surface Area Exposed

Larger surface area = faster total corrosion

• Steel wool (high surface area) rusts in hours
• Steel bar (low surface area) takes days
• Powdered iron rusts almost instantly!

7. Presence of Oxygen

More O₂ = faster corrosion

• Cathode reaction requires O₂
• Areas with restricted O₂ (under paint bubbles, in crevices) can become anodic
• Differential aeration cells form

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Types of Corrosion

1. Uniform Corrosion

Even corrosion over entire surface.

Example: Rusting of unpainted iron sheet in uniform environment

Least dangerous - predictable, easy to design for

2. Pitting Corrosion

Localized attack creating pits/holes.

Most dangerous - hard to detect, causes structural failure

Causes:

• Local breakdown of protective film
• Chloride ions (seawater)
• Differential aeration

Example: Pinholes in stainless steel tanks

3. Galvanic Corrosion

Two dissimilar metals in contact in presence of electrolyte.

More reactive metal (anode) corrodes faster Less reactive metal (cathode) is protected

Example: Steel screws in copper plumbing

• Steel (E° = -0.44 V) corrodes
• Copper (E° = +0.34 V) protected

Prevention: Use similar metals, or insulate dissimilar metals

4. Crevice Corrosion

Corrosion in narrow gaps (crevices, under gaskets, overlapping plates)

Mechanism:

• O₂ depleted in crevice → becomes anodic
• Outside area (high O₂) → cathodic
• Crevice corrodes rapidly!

Example: Under bolt heads, washers, deposits

5. Stress Corrosion

Combination of tensile stress + corrosive environment

Results: Sudden brittle fracture of normally ductile metal

Example: Brass in ammonia atmosphere, stainless steel in chlorides

6. Intergranular Corrosion

Corrosion along grain boundaries of metal crystals

Causes:

• Impurity segregation at grain boundaries
• Depletion of corrosion-resistant elements

Example: Sensitized stainless steel (chromium depleted at grain boundaries)

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Prevention of Corrosion

1. Barrier Protection

Principle: Isolate metal from environment (air, moisture)

(a) Painting/Coating

• Apply paint, lacquer, or polymer coating
• Prevents contact with O₂ and H₂O
• Problem: Scratches expose metal → local corrosion

Examples:

• Car paint (multiple layers!)
• Appliance enamel coating
• Marine paint on ships

(b) Greasing/Oiling

• Thin layer of grease or oil
• Water-repellent barrier
• Used: Machinery, tools, firearms

(c) Electroplating

Deposit a thin layer of corrosion-resistant metal

Two types:

Anodic coating (Sacrificial):

• More reactive metal plated on base
• Coating corrodes preferentially
• Base metal protected even if coating scratched
• Example: Galvanization (Zn on Fe)

$$\text{Zn (E° = -0.76 V)} \rightarrow \text{Zn}^{2+} + 2e^-$$

(Zn corrodes)

$$\text{Fe (E° = -0.44 V)}$$

(protected)

Cathodic coating (Barrier):

• Less reactive metal plated on base
• Acts as barrier only
• Base metal corrodes if coating scratched!
• Example: Tin plating on steel (tin cans), chromium plating

$$\text{Sn (E° = -0.14 V)}$$

(protected)

$$\text{Fe (E° = -0.44 V)} \rightarrow \text{Fe}^{2+} + 2e^-$$

(Fe corrodes if exposed!)

2. Cathodic Protection

Make the entire metal structure the CATHODE of an electrochemical cell.

Since reduction occurs at cathode, metal won’t corrode!

(a) Sacrificial Anode Method

Attach a more reactive metal (lower E°) to the structure.

More reactive metal becomes anode → corrodes Protected structure becomes cathode → safe!

Examples:

Protecting iron ships:

• Attach zinc or magnesium blocks to hull
• Zn (E° = -0.76 V) or Mg (E° = -2.37 V) acts as anode
• Fe (E° = -0.44 V) becomes cathode
• Zn/Mg corrodes, ship protected
• Replace Zn/Mg blocks periodically

Protecting underground pipelines:

• Bury magnesium rods connected to pipeline
• Mg corrodes, pipeline protected
• Used for oil/gas pipelines worldwide

Cell formed:

$$\text{Anode: Mg} \rightarrow \text{Mg}^{2+} + 2e^-$$ $$\text{Cathode: Fe}^{2+} + 2e^- \rightarrow \text{Fe}$$

(reverses corrosion!)

(b) Impressed Current Method

Apply external DC voltage to make structure cathode.

Setup:

• Connect structure to negative terminal of DC source
• Connect inert anode (graphite, platinum) to positive terminal
• Both in same electrolyte (soil, water)

Result:

• Structure forced to be cathode (reduction occurs)
• Inert anode oxidizes (but doesn’t corrode much)
• No metal loss from structure!

Advantages:

• No need to replace sacrificial anodes
• Can control protection level
• Works for large structures

Uses:

• Underground storage tanks
• Marine structures (offshore platforms)
• Large pipelines

3. Passivation

Form a thin, protective oxide layer on metal surface.

Mechanism:

• Oxide layer is non-porous and adherent
• Prevents further oxidation
• Self-healing if scratched (in presence of O₂)

Examples:

Aluminum:

$$4\text{Al} + 3\text{O}_2 \rightarrow 2\text{Al}_2\text{O}_3$$

• Forms instantly in air
• Al₂O₃ layer is only 4-6 nm thick
• Prevents further corrosion
• Can be thickened by anodization

Stainless steel:

• Contains Cr (>12%)
• Forms Cr₂O₃ protective layer
• Highly corrosion-resistant
• Used in cutlery, medical instruments

Titanium:

• Forms TiO₂ layer
• Extremely protective
• Used in aerospace, medical implants

4. Alloying

Add elements to improve corrosion resistance

Examples:

Stainless steel:

• Fe + 18% Cr + 8% Ni + small amounts of C, Mn
• Chromium forms protective oxide
• Nickel improves ductility and corrosion resistance
• Used in utensils, construction, chemical plants

Brass (Cu + Zn):

• More corrosion-resistant than pure copper in some environments
• Used in marine hardware, plumbing

Bronze (Cu + Sn):

• Excellent corrosion resistance
• Ancient bronze statues still intact!
• Ship propellers, bells, sculptures

5. Environmental Modification

Change the corrosive environment

Methods:

Dehumidification:

• Remove moisture from air
• Use silica gel, calcium chloride desiccants
• Nitrogen purging in storage

pH control:

• Add inhibitors to maintain neutral/alkaline pH
• Used in cooling water systems

Oxygen removal:

• Deaerate water (boil, vacuum)
• Add oxygen scavengers (sodium sulfite)
• Used in boiler water treatment

Inhibitors:

• Add chemicals that form protective films
• Example: Chromates, phosphates in cooling systems

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Practice Problems

Level 1: JEE Main

Q1. Which of the following metals does NOT corrode in moist air?

• (a) Iron (b) Aluminum (c) Gold (d) Zinc

Q2. In galvanized iron, which metal acts as a sacrificial anode?

• (a) Iron (b) Zinc (c) Tin (d) Chromium

Q3. What are the essential requirements for rusting of iron?

Q4. Write the anodic and cathodic reactions during rusting of iron.

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Level 2: JEE Main/Advanced

Q5. Explain why tin-plated iron corrodes more rapidly than galvanized iron when the coating is broken.

Q6. A ship’s hull is protected by attaching magnesium blocks. If 10 kg of Mg is consumed in 1 year, how much iron was saved from corrosion? Given: M(Mg) = 24, M(Fe) = 56

Q7. Arrange the following in increasing order of corrosion rate: Pure iron, iron with carbon impurities, stainless steel, gold

Q8. Why does aluminum not corrode rapidly despite being more reactive than iron?

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Level 3: JEE Advanced

Q9. In cathodic protection of an iron pipeline using impressed current method:

• (a) What is connected to the negative terminal?
• (b) What happens at the iron surface?
• (c) Why doesn’t the iron corrode?

Q10. A steel structure is protected using zinc sacrificial anodes. The steel structure has surface area 100 m² and average corrosion rate would be 0.1 mm/year without protection. If all corrosion is transferred to zinc, calculate the mass of zinc consumed per year. Given: Density of Fe = 7.86 g/cm³, M(Fe) = 56, M(Zn) = 65

Q11. Explain the mechanism of differential aeration corrosion. Why does iron corrode faster under a water drop?

Q12. Compare the effectiveness of the following for protecting iron:

• (a) Tin coating
• (b) Zinc coating
• (c) Chromium plating
• (d) Cathodic protection

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Solutions to Practice Problems

A1. (c) Gold - E° = +1.50 V (very unreactive, doesn’t oxidize)

A2. (b) Zinc - More reactive than iron (E° = -0.76 V)

A3.

1. Oxygen (O₂)
2. Water or moisture
3. Electrolyte (optional but accelerates)

A4. Anode: Fe → Fe²⁺ + 2e⁻ Cathode: O₂ + 2H₂O + 4e⁻ → 4OH⁻

A5. Tin-plated: Sn (E° = -0.14 V) is less reactive than Fe (E° = -0.44 V). When coating breaks, Fe becomes anode and corrodes rapidly with Sn as cathode.

Galvanized: Zn (E° = -0.76 V) is more reactive than Fe. Even when broken, Zn corrodes preferentially, protecting Fe.

A6. Mg²⁺ + 2e⁻ → Mg (n=2), Fe²⁺ + 2e⁻ → Fe (n=2) Moles of Mg = 10,000/24 = 416.67 mol Same moles of Fe saved = 416.67 mol Mass of Fe = 416.67 × 56 = 23,333 g = 23.3 kg

A7. Gold < Stainless steel < Pure iron < Iron with carbon impurities

A8. Aluminum forms a thin, adherent, non-porous Al₂O₃ layer (passivation) that protects the underlying metal from further oxidation.

A9. (a) Iron pipeline (cathode) (b) Reduction occurs: Fe²⁺ + 2e⁻ → Fe (reverses corrosion) (c) Iron is forced to be cathode where only reduction occurs, preventing oxidation (corrosion)

A10. Volume of Fe that would corrode = 100 m² × 0.0001 m = 0.01 m³ = 10,000 cm³ Mass of Fe = 10,000 × 7.86 = 78,600 g Moles of Fe = 78,600/56 = 1403.6 mol Fe²⁺ + 2e⁻ → Fe (n=2), Zn²⁺ + 2e⁻ → Zn (n=2) Same electrons, so same moles: 1403.6 mol Zn Mass of Zn = 1403.6 × 65 = 91,233 g = 91.2 kg

A11. Under water drop:

• Center (low O₂) → anode → Fe → Fe²⁺ + 2e⁻
• Edge (high O₂) → cathode → O₂ + 2H₂O + 4e⁻ → 4OH⁻
• Differential O₂ concentration creates galvanic cell
• Iron corrodes faster at center (pit forms)

A12. Effectiveness ranking (best to worst):

1. Cathodic protection - Active prevention, iron can’t oxidize
2. Zinc coating - Sacrificial, protects even when scratched
3. Chromium plating - Good barrier, some protection if thin
4. Tin coating - Only barrier, accelerates corrosion if broken

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Common Mistakes to Avoid

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Key Takeaways for JEE

Essential Concepts

1. Corrosion = spontaneous electrochemical oxidation of metals
2. Rusting requires: O₂ + H₂O (+ electrolyte accelerates)
3. Anode reaction: Fe → Fe²⁺ + 2e⁻ (iron corrodes)
4. Cathode reaction: O₂ + 2H₂O + 4e⁻ → 4OH⁻
5. Sacrificial protection: More reactive metal corrodes, protects base metal
6. Cathodic protection: Force structure to be cathode (no oxidation)

Must Remember

Protection Method	Principle	Example
Galvanization	Zn coating (sacrificial)	Iron sheets, pipes
Tin plating	Sn barrier (not sacrificial!)	Tin cans
Cathodic protection	Make structure cathode	Ships, pipelines
Passivation	Protective oxide layer	Aluminum, stainless steel
Painting	Barrier coating	Cars, buildings
Alloying	Add corrosion-resistant elements	Stainless steel

Electrode Potential Ladder

$$\text{Mg (-2.37 V)} < \text{Zn (-0.76 V)} < \text{Fe (-0.44 V)} < \text{Sn (-0.14 V)} < \text{Cu (+0.34 V)}$$

More negative → Corrodes first → Good for sacrificial protection

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Real-World Applications

Modern Examples (2025)

Golden Gate Bridge:

• Constant repainting (every 7 years)
• 10,000 gallons of paint used
• Protects 1.2 million rivets from corrosion
• Cost: $40 million per cycle

Mars Rovers:

• Use gold-plated contacts (no oxidation on Mars!)
• Titanium bodies (passive oxide layer)
• Still working after 15+ years

Oil Rigs:

• Impressed current cathodic protection
• Sacrificial aluminum anodes
• Coatings resistant to seawater
• Lifetime: 25-30 years with proper protection

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Related Topics

Within Electrochemistry

• Oxidation-Reduction — Electron transfer in corrosion
• Electrochemical Cells — Corrosion as galvanic cell
• Electrode Potentials — Predicting which metal corrodes
• Electrolysis — Cathodic protection mechanism

Cross-Chapter Connections

• Metallurgy — Metal extraction and protection
• Chemical Kinetics — Rate of corrosion
• Thermodynamics — Spontaneity of corrosion

Physics Connections

• Current Electricity — Electron flow in corrosion cells
• Electrostatics — Potential differences

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