Quick Notes

Quick Notes

Chemistry Chapter 1: Redox Equilibrium
1.6 Rusting
1.6 Rusting - Chemistry Form 5
CHEMISTRY • Form 5 • Chapter 1: Redox Equilibrium

1.6Rusting

Learn the redox mechanism of iron rusting, the effects of other metals on corrosion, and methods to prevent rusting.

Learning Objectives

  • Describe metal corrosion process as redox reaction through activities.
  • Experiment to prevent rusting.

Electrochemical Mechanism of Iron Rusting in a Water Droplet

An isolated water droplet on an iron surface acts as a complete simple chemical cell.

Brief Explanation

Understand the key ideas behind metal corrosion and rusting.

Metal corrosion is a redox reaction where a metal is oxidised spontaneously when metal atoms release electrons to form metal ions: \[M \rightarrow M^{n+} + ne^-\] The more electropositive the metal is, the easier it is for the metal to corrode.

1. Rusting of Iron

Rusting is a chemical corrosion process that occurs when iron is exposed to water and oxygen, undergoing a redox reaction.

Rust is hydrated iron(III) oxide, \(\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}\), which is reddish-brown, brittle, cracked, and permeable. It does not stick strongly to the metal surface, allowing oxygen and water to permeate continuously and damage the underlying iron.

2. Non-porous Protective Oxide Layers

Unlike iron, metals like aluminium form a hard, non-porous layer of aluminium oxide that protects underlying aluminium from corroding.

Other metals exhibiting similar protective oxide layers include zinc, lead, nickel, and chromium.

Experiment 1D: Corrosion of Copper and Iron Details
Aim To study the corrosion of copper and iron.
Hypothesis Corrosion of copper metal produces blue precipitate and corrosion of iron metal produces brown precipitate.
Manipulated variable Copper and iron.
Responding variable Presence of blue and brown precipitate.
Fixed variable Salt solution.
Reagent roles Sodium hydroxide, \(\text{NaOH}\), provides \(\text{OH}^-\) ions to precipitate metal ions as coloured hydroxides. \(\text{Cu}^{2+}\) forms blue copper(II) hydroxide, \(\text{Cu(OH)}_2\). \(\text{Fe}^{2+}\) forms green iron(II) hydroxide, \(\text{Fe(OH)}_2\), which oxidises to reddish-brown iron(III) hydroxide / hydrated iron(III) oxide.

Mechanism of Iron Rusting

When an iron surface comes into contact with a water droplet and oxygen, a simple chemical cell is established.

  • Anode (Negative terminal) — Oxidation: Located at the surface of iron in the middle of the water droplet where oxygen concentration is low. Iron atom loses electrons: \(\text{Fe}(s) \rightarrow \text{Fe}^{2+}(aq) + 2e^-\).
  • Electron conduction: Released electrons flow through the iron to the outer edge of the water droplet where oxygen concentration is higher.
  • Cathode (Positive terminal) — Reduction: Located at the iron surface at the edge of the droplet where oxygen concentration is high. Dissolved oxygen accepts electrons and is reduced to hydroxide ions: \(\text{O}_2(g) + 2\text{H}_2\text{O}(l) + 4e^- \rightarrow 4\text{OH}^-(aq)\).
  • Precipitation and Further Oxidation: \(\text{Fe}^{2+}\) ions combine with \(\text{OH}^-\) ions to form green iron(II) hydroxide precipitate: \(\text{Fe}^{2+}(aq) + 2\text{OH}^-(aq) \rightarrow \text{Fe(OH)}_2(s)\). Iron(II) hydroxide undergoes continuous oxidation by dissolved oxygen to form hydrated iron(III) oxide (rust), where \(x\) is an integer of variable values: \(\text{Fe(OH)}_2(s) \xrightarrow{\text{oxidation}} \text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}(s)\).

Electrolyte Effects

Rusting occurs faster in the presence of acids or salts because dissolved ions turn water into a better electrolyte, increasing electrical conductivity.

Experiment 1E: Effects of Other Metals on Rusting Details
Aim To study the effects of other metals on rusting.
Hypothesis When a more electropositive metal comes into contact with an iron nail, the iron does not rust; when a less electropositive metal comes into contact with an iron nail, the iron nail rusts.
Manipulated variable Different metals coiled around the iron nail.
Responding variable Rusting of iron.
Fixed variable Iron nail.
Indicators and reagents Potassium hexacyanoferrate(III), \(\text{K}_3\text{Fe(CN)}_6\), detects \(\text{Fe}^{2+}\) ions by forming an intense dark blue precipitate/colour. Phenolphthalein turns pink in the presence of \(\text{OH}^-\) ions produced at cathode regions. Hot agar solution is used instead of water to slow down convection/diffusion, keeping coloured spots and precipitates immobilised for clear observation.
Control An iron nail without any coiled metal coil (test tube T).

More Electropositive Metals (Mg, Al, Zn)

Slow down or prevent rusting. They release electrons more easily than iron (\(\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-\)). The coiled metal corrodes (sacrificial protection), while reduction occurs on the iron surface (\(\text{O}_2 + 2\text{H}_2\text{O} + 4e^- \rightarrow 4\text{OH}^-\)), showing pink colour and no blue colour.

Less Electropositive Metals (Sn, Pb, Cu)

Speed up rusting of iron. Iron is more electropositive than these metals, so iron atoms lose electrons more readily (\(\text{Fe} \rightarrow \text{Fe}^{2+} + 2e^-\)), forming high concentrations of \(\text{Fe}^{2+}\) (intense dark blue precipitate).

Order of tendency to release electrons (electropositivity): \[\text{Mg} > \text{Al} > \text{Zn} > \text{Fe} > \text{Sn} > \text{Pb} > \text{Cu}\]

1. Use of Protective Surface

A barrier against oxygen and water.

  • Paint, grease, and plastic wrappings: Paint on fences, cars, and bridges; oil/grease on car engines; plastic coating on wire fences.
  • Galvanisation: Coating iron or steel with a thin zinc layer. Zinc forms a tough, non-porous zinc oxide layer. If scratched, zinc still protects iron via sacrificial protection because zinc is more electropositive than iron.
  • Tin plating: Plating steel plates with a thin tin layer (used in food cans). Tin forms a protective oxide layer. If a tin-plated can is dented or scratched, iron underneath is exposed to air and water; because iron is more electropositive than tin, iron rusts even faster.
  • Chrome plating: Used on car bumpers, bicycles, and decorative items.

2. Sacrificial Protection

Iron is attached or connected to a more electropositive sacrificial metal (e.g., magnesium or zinc). The sacrificial metal corrodes preferentially and must be replaced periodically.

Applications: Zinc blocks/plates on ship hulls; magnesium blocks connected to underground steel pipes or offshore oil rig pillars.

3. Alloying

Iron is alloyed with carbon, chromium, and nickel to form stainless steel. Chromium and nickel form a shiny, protective, and corrosion-proof oxide layer that shields the iron underneath from oxygen and water.

Applications: Kitchen utensils and surgical instruments.

Try to Answer First

Answer mentally, then check the explanation.

1

Why does iron rust faster when exposed to seawater compared to pure rainwater?

Answer: Seawater contains dissolved salts (such as sodium chloride) which act as electrolytes, significantly increasing the electrical conductivity of water and accelerating electron transfer during the redox reaction.
2

What is the function of potassium hexacyanoferrate(III) solution in experiments studying iron corrosion?

Answer: It acts as a detecting reagent for iron(II) ions, \(\text{Fe}^{2+}\); it forms an intense dark blue precipitate in the presence of \(\text{Fe}^{2+}\), indicating the occurrence and extent of iron rusting.
3

Why is it dangerous to consume food from a dented tin can?

Answer: When a tin-plated can is dented, the protective tin coating may crack, exposing the steel (iron) underneath. Because iron is more electropositive than tin, iron loses electrons faster and corrodes rapidly in contact with tin, causing the food to become contaminated with rust.

Common Mistakes

Avoid these errors when studying rusting.

!
Writing the reduction half-equation at the cathode incorrectly (e.g., writing the reduction of \(\text{H}^+\) or \(\text{Fe}^{2+}\) instead of neutral dissolved oxygen: \(\text{O}_2 + 2\text{H}_2\text{O} + 4e^- \rightarrow 4\text{OH}^-\)).
!
Stating that metals more electropositive than iron accelerate rusting; in fact, more electropositive metals slow down or prevent rusting by acting as sacrificial anodes.
!
Forgetting that rust requires both water and oxygen; absence of either prevents rusting.
!
Confusing phenolphthalein (detects \(\text{OH}^-\) ions, turning pink) with potassium hexacyanoferrate(III) (detects \(\text{Fe}^{2+}\) ions, turning dark blue).

Concept Misunderstandings

Compare each misconception with the correct fact.

Misconception

All metal corrosion produces brittle rust like iron.

Correct Fact

Rusting refers strictly to the corrosion of iron. Other metals like aluminium, zinc, and chromium corrode to form hard, non-porous, protective oxide coatings that stop further corrosion.

Misconception

Tin plating prevents rusting by sacrificial protection.

Correct Fact

Tin plating acts only as a protective barrier surface. Tin is less electropositive than iron, so if the tin layer is damaged, iron corrodes even more rapidly. Sacrificial protection requires a more electropositive metal like zinc or magnesium.

Summary

Key facts to remember.

  • Corrosion of metal is a spontaneous oxidation process (\(M \rightarrow M^{n+} + ne^-\)).
  • Rusting of iron forms hydrated iron(III) oxide, \(\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}\), requiring both \(\text{O}_2\) and \(\text{H}_2\text{O}\).
  • Rusting mechanism: Anode in low-\(\text{O}_2\) area (\(\text{Fe} \rightarrow \text{Fe}^{2+} + 2e^-\)); Cathode in high-\(\text{O}_2\) area (\(\text{O}_2 + 2\text{H}_2\text{O} + 4e^- \rightarrow 4\text{OH}^-\)); Precipitation (\(\text{Fe}^{2+} + 2\text{OH}^- \rightarrow \text{Fe(OH)}_2\)) followed by oxidation to rust.
  • Contact with more electropositive metals (\(\text{Mg}, \text{Al}, \text{Zn}\)) prevents rusting (sacrificial protection). Contact with less electropositive metals (\(\text{Sn}, \text{Pb}, \text{Cu}\)) accelerates rusting.
  • Prevention methods: Protective barrier surfaces (paint, oil, grease, plastic, tin/chrome plating, galvanisation), sacrificial protection (zinc/magnesium blocks), and alloying (stainless steel).

Short Activity

Complete all parts before checking your score.

A. Objective Quiz

1

In the mechanism of iron rusting, which half-equation takes place at the positive terminal (cathode)?

2

An iron nail is coiled with a strip of metal and placed in agar containing potassium hexacyanoferrate(III) solution. Which metal will cause an intense dark blue precipitate to appear around the iron nail most rapidly?

B. Fill in the Blanks

3

In sacrificial protection of underground pipes and ship hulls, iron acts as the cathode while a more electropositive metal such as ________ acts as the sacrificial anode by undergoing oxidation.

4

In the corrosion detection experiment, ________ indicator is used to detect the presence of hydroxide ions, \(\text{OH}^-\), by producing a pink coloration.

C. Match the Prevention Method with Its Description

Drag each method on desktop. On mobile, tap a method and then tap the matching box.

Prevention Method

Galvanisation Sacrificial protection Alloying Tin plating

Description / Application

1 Iron is coated with zinc which forms a non-porous oxide and provides sacrificial protection even if scratched.
2 Connecting magnesium or zinc blocks to ship hulls and steel pillars of oil rigs.
3 Iron is combined with carbon, chromium, and nickel to make stainless steel for surgical instruments.
4 Used on food cans to form an oxide barrier, but accelerates rusting if the surface is scratched.
Complete every answer before checking.
 

Keywords

Tap or hover over a term to read its meaning.

Corrosion of metal Rusting Rust Sacrificial protection Galvanisation Potassium hexacyanoferrate(III) Alloying