Corrosion of metals Revision Notes

Rusting is a specific type of corrosion that affects iron and steel, leading to the formation of a flaky, reddish-brown solid called rust. Rusting weakens the metal by converting it into a brittle and porous material, which can eventually cause structural failure if untreated.

Rusting requires the presence of both oxygen and water. Neither oxygen alone nor water alone is sufficient to cause rusting rapidly; both must be present simultaneously for the process to proceed.

• Rusting is the corrosion of iron and steel.
• It produces hydrated iron(III) oxide, Fe2O3·xH2O, known as rust.
• Both oxygen and water must be present for rusting to occur.

The two essential conditions for rusting are: exposure to oxygen (from air) and the presence of water (moisture). Water acts as an electrolyte, facilitating the electrochemical reactions that cause rusting.

In dry air, iron does not rust because water is absent. Similarly, in pure water without dissolved oxygen, rusting is very slow or negligible.

The presence of dissolved salts in water increases its conductivity, accelerating rusting.

• Oxygen from air must be present.
• Water (moisture) must be present to act as an electrolyte.
• Salt or impurities in water increase rusting rate by enhancing conductivity.

Rusting is an electrochemical process involving oxidation and reduction reactions occurring at different sites on the iron surface.

At the anodic sites, iron atoms lose electrons (oxidation) to form iron(II) ions:

Fe(s) → Fe²⁺(aq) + 2e⁻ (oxidation) At the cathodic sites, oxygen from air gains electrons (reduction) in the presence of water to form hydroxide ions: O₂(g) + 4e⁻ + 2H₂O(l) → 4OH⁻(aq) (reduction) The iron(II) ions react with hydroxide ions to form iron(II) hydroxide, which is further oxidised by oxygen and water to form hydrated iron(III) oxide (rust). The overall simplified reaction is: 4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃ Iron(III) hydroxide then dehydrates to form hydrated iron(III) oxide, Fe₂O₃·xH₂O, the reddish-brown rust.

• Anodic reaction: Fe(s) → Fe²⁺(aq) + 2e⁻
• Cathodic reaction: O₂ + 4e⁻ + 2H₂O → 4OH⁻
• Overall reaction: 4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃
• Rust is hydrated iron(III) oxide, Fe₂O₃·xH₂O.

Rust is a hydrated iron(III) oxide with the general formula Fe₂O₃·xH₂O, where x represents variable amounts of water molecules incorporated in the solid.

It appears as a reddish-brown, flaky, porous solid that does not adhere well to the underlying metal, allowing further corrosion to continue beneath the rust layer.

The formation of rust weakens iron and steel by converting strong metallic bonds into brittle oxide layers.

• Rust formula: Fe₂O₃·xH₂O (hydrated iron(III) oxide)
• Rust is reddish-brown and flaky.
• Rust does not protect iron from further corrosion.

Write a balanced chemical equation for the rusting of iron in the presence of oxygen and water to form hydrated iron(III) oxide.

Step 1: Write the reactants and products: iron (Fe), oxygen (O₂), water (H₂O), and hydrated iron(III) oxide (Fe₂O₃·xH₂O).

Step 2: Balance iron atoms: 4Fe on the left to match 4Fe on the right (as 2 Fe₂O₃ units). Step 3: Balance oxygen atoms: 3O₂ molecules provide 6 oxygen atoms, matching 6 oxygen atoms in 4Fe(OH)₃ (which can be considered as hydrated iron(III) oxide). Step 4: Balance hydrogen atoms: 6H₂O molecules provide 12 hydrogen atoms, matching 12 hydrogen atoms in 4Fe(OH)₃. Step 5: Write the balanced equation: 4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃ This equation represents the formation of iron(III) hydroxide, which dehydrates to form hydrated iron(III) oxide (rust).

• Balanced equation: 4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃
• Fe(OH)₃ dehydrates to Fe₂O₃·xH₂O (rust).

Rusting is a chemical process where iron or steel reacts with oxygen and water to form hydrated iron(III) oxide, commonly known as rust. This reaction weakens the metal, causing it to flake and lose strength. Preventing rust involves stopping oxygen and water from reaching the metal surface. Barrier methods achieve this by creating a protective layer that isolates the metal from the environment.

• Rusting requires oxygen and water.
• Barrier methods block oxygen and water from contacting metal.

Painting involves applying a layer of paint to the metal surface. The paint forms a continuous film that physically blocks air and moisture from reaching the metal. This method is widely used because it is relatively inexpensive and can be applied to large surfaces.

Paint must be applied evenly and maintained to avoid cracks or chips, which would expose the metal and allow rusting to start.

• Paint forms a protective film over metal.
• Prevents oxygen and water contact.
• Requires maintenance to avoid damage.

Greasing involves coating the metal surface with a layer of grease or oil. The greasy layer repels water and reduces oxygen diffusion to the metal surface. Greasing is commonly used for moving parts, such as hinges or machinery components, where paint might crack due to movement.

Grease also reduces friction and wear, providing dual benefits. However, it can attract dust and dirt, so regular cleaning and reapplication are necessary.

• Grease repels water and oxygen.
• Suitable for moving parts.
• Needs regular cleaning and reapplication.

Plastic coatings, such as PVC or other polymer films, can be applied to metal surfaces to create a waterproof and airtight barrier. These coatings are durable and resistant to environmental damage, making them effective for long-term protection.

Plastic coatings are often used on pipes, outdoor furniture, and tools. The coating must be intact and free from punctures to maintain protection.

• Plastic coatings block moisture and air.
• Durable and long-lasting protection.
• Must be free from damage to be effective.

Rusting is the corrosion of iron that occurs when iron reacts with oxygen and water. The overall chemical process produces hydrated iron(III) oxide, commonly known as rust. The presence of both oxygen and water is essential for rusting to occur.

The simplified overall reaction for rusting can be represented as:

4Fe(s) + 3O2(g) + 6H2O(l) → 4Fe(OH)3(s) (hydrated iron(III) oxide)

• Rusting requires iron, oxygen, and water.
• Rust is hydrated iron(III) oxide, Fe2O3·xH2O.
• Both oxygen and water must be present for rusting to proceed.

Barrier methods work by physically preventing oxygen and water from reaching the iron surface. Without access to these reactants, the electrochemical reactions that cause rusting cannot occur.

These methods create a continuous, impermeable layer over the iron, isolating it from the environment. This stops the oxidation of iron atoms and the subsequent formation of rust.

If the barrier is damaged or incomplete, oxygen and water can penetrate and rusting will begin at those points.

• Barrier methods block oxygen and water from contacting iron.
• They prevent the electrochemical reactions that cause rusting.
• Integrity of the barrier is essential for effective protection.

1. Painting: Applying paint forms a solid coating that seals the iron surface from air and moisture. Paint must be applied evenly and maintained to avoid cracks or chips.

2. Oiling or Greasing: A layer of oil or grease repels water and reduces oxygen contact. This method is often used for tools or machinery parts.

3. Plastic Coatings: Plastic films or coatings provide a waterproof and airtight layer over iron surfaces, preventing rusting.

• Painting creates a solid, impermeable layer.
• Oiling repels water and limits oxygen access.
• Plastic coatings provide a durable barrier against moisture and air.

Galvanizing involves coating iron or steel with a thin layer of zinc. Zinc acts as a physical barrier preventing oxygen and water from reaching the iron surface.

The zinc layer adheres tightly and protects the iron underneath. If the zinc layer is scratched, it still offers some protection by corroding preferentially (sacrificial protection), but this is beyond the scope of barrier methods.

Galvanizing is widely used for outdoor steel structures due to its durability.

• Zinc coating forms a protective barrier over iron.
• Prevents oxygen and water from reaching iron surface.
• Commonly used for outdoor protection of steel.

The effectiveness of barrier methods depends on the completeness and durability of the protective layer. Any breaks or damage allow oxygen and water to contact iron, initiating rusting.

Regular maintenance, such as repainting or re-oiling, is necessary to maintain protection.

Barrier methods do not stop rusting chemically but prevent the reactants from reaching iron, thus stopping the rusting process physically.

• Complete, undamaged barriers are essential.
• Maintenance extends the life of the barrier.
• Barrier methods prevent rusting by physical exclusion of oxygen and water.

Corrosion is the chemical deterioration of metals caused by reactions with substances in the environment, especially oxygen and water. Iron and steel commonly rust, forming hydrated iron(III) oxide (Fe2O3·xH2O), which weakens the metal.

Galvanizing is a process where a protective layer of zinc is applied to iron or steel surfaces to prevent corrosion. This method utilises zinc’s chemical properties to protect the underlying metal by two main mechanisms: barrier protection and sacrificial protection.

• Corrosion involves oxidation of iron to iron ions.
• Rusting requires oxygen and water.
• Galvanizing coats metal with zinc to prevent rust.

In galvanizing, zinc forms a continuous coating over the iron or steel surface. This physical layer acts as a barrier, preventing oxygen and moisture from reaching the metal beneath. Without access to these reactants, the iron cannot oxidise and rust.

The barrier is effective only if the zinc coating is intact and free from cracks or scratches. Any damage exposes the iron surface, allowing corrosion to start unless sacrificial protection intervenes.

• Zinc coating blocks oxygen and water.
• Prevents direct contact between iron and corrosive agents.
• Effectiveness depends on coating integrity.

Zinc is more reactive than iron, meaning it oxidises more readily. If the zinc coating is scratched or damaged, exposing the iron, zinc will corrode preferentially. This is called sacrificial protection because zinc 'sacrifices' itself to protect the iron.

The chemical basis involves zinc atoms losing electrons (oxidation) to form zinc ions, while the iron remains unoxidised. This process continues until the zinc is consumed, after which the iron may begin to corrode if exposed.

The relevant half-equations are: Zn(s) → Zn²⁺(aq) + 2e⁻ (oxidation) Fe(s) → Fe²⁺(aq) + 2e⁻ (oxidation, less likely when zinc present) The electrons released by zinc oxidation prevent iron atoms from losing electrons, thus protecting iron from rusting.

• Zinc corrodes before iron due to higher reactivity.
• Zinc oxidation protects iron from oxidising.
• Sacrificial protection continues until zinc is used up.

When zinc corrodes, it reacts with oxygen and water to form zinc hydroxide, which can further react to form zinc oxide or other compounds, depending on conditions.

A simplified overall reaction for zinc corrosion in moist air is: 2Zn(s) + O2(g) + 2H2O(l) → 2Zn(OH)2(s) Zinc hydroxide may dehydrate to zinc oxide (ZnO), which forms a protective layer slowing further corrosion.

In contrast, iron exposed to the same conditions forms hydrated iron(III) oxide (rust), which flakes off and exposes fresh metal, accelerating corrosion. Zinc corrosion products adhere better, maintaining protection.

• Zinc reacts with oxygen and water to form zinc hydroxide.
• Zinc corrosion products form a protective layer.
• Iron corrosion products (rust) are flaky and non-protective.

Galvanizing extends the life of iron and steel structures such as bridges, fences, and roofing by preventing rust.

The dual protection mechanisms mean that even if the zinc coating is scratched, the underlying metal remains protected for a significant time.

Galvanized metals require less maintenance and replacement, reducing costs and improving safety in construction and manufacturing.

• Galvanizing increases metal lifespan.
• Sacrificial protection ensures durability despite coating damage.
• Widely used in construction and industry.

Sacrificial protection is a method used to prevent the corrosion of a metal by attaching it to a more reactive metal. The more reactive metal acts as a sacrificial anode and corrodes instead of the metal being protected. This method is commonly used to protect iron or steel structures from rusting.

Corrosion involves the oxidation of metals, where metal atoms lose electrons to form positive ions. By using a more reactive metal that oxidises more readily, the less reactive metal remains protected from corrosion.

• Sacrificial protection prevents corrosion by using a more reactive metal.
• The more reactive metal corrodes (oxidises) instead of the protected metal.
• Commonly used to protect iron and steel from rusting.

The reactivity series is a list of metals arranged in order of their reactivity, from the most reactive to the least reactive. Metals higher in the series lose electrons more easily and oxidise more readily.

In sacrificial protection, the metal higher in the reactivity series acts as the sacrificial anode. For example, zinc is more reactive than iron, so zinc will oxidise first when both metals are in contact.

This difference in reactivity causes the more reactive metal to lose electrons (oxidation), protecting the less reactive metal from corrosion.

• Reactivity series ranks metals by ease of electron loss.
• More reactive metals oxidise before less reactive metals.
• Zinc (more reactive) protects iron (less reactive) by corroding first.

Corrosion is an oxidation process where metal atoms lose electrons to form metal ions. In sacrificial protection, the more reactive metal loses electrons (oxidation) and forms positive ions.

These electrons flow to the less reactive metal, which gains electrons and is thus reduced or prevented from oxidising.

For example, when zinc protects iron, zinc atoms lose electrons to form Zn²⁺ ions, while iron atoms do not lose electrons and remain unoxidised.

• Oxidation is loss of electrons; reduction is gain of electrons.
• Sacrificial anode metal loses electrons and corrodes.
• Electrons flow to protected metal, preventing its oxidation.

Consider zinc protecting iron in the presence of water and oxygen, conditions that normally cause iron to rust.

Zinc oxidises by losing electrons: Zn(s) → Zn²⁺(aq) + 2e⁻

The electrons released flow to the iron, preventing iron atoms from losing electrons and forming Fe²⁺ or Fe³⁺ ions, thus preventing rusting (iron oxidation). The overall effect is that zinc corrodes instead of iron.

• Zn(s) → Zn²⁺(aq) + 2e⁻ (oxidation of zinc)
• Electrons protect iron from oxidation.
• Iron remains unoxidised and does not rust.

Sacrificial protection is used in galvanising, where a coating of zinc is applied to iron or steel. The zinc layer acts as a sacrificial anode.

Even if the zinc coating is scratched, the exposed iron remains protected because zinc corrodes preferentially.

This method requires the two metals to be in electrical contact and exposed to an electrolyte (e.g., water containing dissolved oxygen) to allow electron flow and ion movement.

• Galvanising applies zinc coating to iron or steel.
• Zinc corrodes preferentially, protecting iron even if scratched.
• Electrical contact and electrolyte presence are necessary for protection.