Corrosion Chemistry Tutorial

Key Concepts

• Corrosion: the oxidation of metals by certain substances such as water and oxygen in the environment.
• Rusting of iron is a common form of corrosion.
• Noble Metal: a metal resistant to corrosion and found uncombined in nature, eg, gold and silver.

  Noble metals are also known as Cathodic Metals, or Protected Metals.

• Ignoble Metal: a metal that corrodes.

  Ignoble metals are also known as Anodic Metals, or Corroding Metals.

• Passivating metal: a reactive metal that forms an inactive coating as a result of reacting with substances such as oxygen and water, eg, aluminium and chromium.
• In general, corrosion is a spontaneous electrochemical process¹ which requires:

  (i) an anode where oxidation occurs

  (ii) a cathode where reduction occurs

  (iii) a metal path to allow for the flow of electrons

  (iv) an electrolyte

• Common methods of corrosion prevention are:

  (i) Painting

  (ii) Plating with another metal

  (iii) Cathodic protection

  (iv) Alloying with another metal

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Example of Corrosion: Rusting of Iron

If you look very closely at a piece of iron as it begins to rust you might notice two things:

• In some places there are flakes of red rust.
• In other places there are small holes, or pits.

This is because the site where you see the red rust is not the site where the iron is initially being oxidised.

In order for the rust to form on the iron, the iron needs to contain an impurity, that is, a place where the electrons released from the oxidation of iron can travel to.

In order to complete the electrical circuit, ions need to be able to travel through an electrolyte, and this is provided by moisture, water in the air surrounding the iron, in which ions can dissolve and travel to the site of rust formation.

The process by which rust is formed on iron is shown in the diagram below:

Fe_(s) is oxidised to Fe²⁺ and a pit develops on the surface of the iron.

Fe_(s) → Fe²⁺_(aq) + 2e^-

At the site of an impurity oxygen is reduced to hydroxide ions.

O_2(g) + 2H₂O_(l) + 4e^- → 4OH^-_(aq)

Electrons flow through the iron bar from the pit to the impurity.

Ions flow through the water electrolyte to the rusting site.

At the rusting site, iron(II) hydroxide, Fe(OH)_2(s), is formed which readily undergoes oxidation to form hydrated iron(III) oxide, Fe₂O₃.H₂O_(s), (rust).

The equations representing the reactions at the anode, cathode, and at the site where the rust forms are shown below:

Factors Influencing the Rate of Corrosion of Iron

The rate of corrosion of iron, or the rate of rusting of iron, refers to how quickly the iron corrodes (or rusts).
A number of factors can influence how fast the iron will rust:

• Purity of the iron:

  Pure iron rusts very, very slowly.
  Rusting of iron occurs very quickly if a less active metal impurity is present.
  For example, copper and tin are less active metals than iron. If copper or tin impurities are present in the iron, the iron will rust very quickly.

• Amount of Stress:

  Bending, cutting and sharpening the iron place it under stress which distorts the arrangement of metal ions in the lattice.
  Fe²⁺ ions can escape from the metal lattice more easily, so rusting occurs faster.

• Electrolyte Composition:

  Pure water is a poor electrical conductor so rusting occurs more slowly.
  Salt water is a much better electrical conductor so rusting occurs more rapidly.

• Oxygen Concentration:

  A large concentration of oxygen at the surface of iron will speed up the corrosion of another area with lower oxygen concentration if there is a path for electrons and ions to follow between the two areas.

Rust Prevention Methods

• Painting

  Prevents oxygen and water coming into contact with the iron so rust can't form.
  If the paint is scratched or chipped, oxygen and water can come into contact with the iron so rust will then form.

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• Plating with Another Metal

  (i) Plating with a less active metal like tin is just like painting. The tin covering prevents the iron coming into contact with water and oxygen, but, if the tin plate is scratched, rust will readily form.

  (ii) Plating with a more active metal like zinc in the galvanising process, firstly acts like painting in that water and oxygen can not get to the iron surface, but, if the surface is scratched this sets up a new galvanic electrochemical cell:

  Zn_(s) + Fe²⁺_(aq) → Zn²⁺_(aq) + Fe_(s)

  The iron is reformed in this process.
  Zinc ions will form zinc oxide (or carbonate) which is less permeable to water and oxygen than iron oxides so this slows down the rate of corrosion.

• Cathodic Protection:

  (i) A Sacrifical Anode is a block of active metal such as zinc or magnesium attached to the hull of a ship, or to tanks and pipelines buried in moist earth.
  The more active zinc sacrifices itself by corroding in preference to the iron.
  Sacrifical anodes must be replaced when they have corroded in order for the protection of iron to be maintained.

  (ii) An Impressed Current System is used to protect immersed hulls of ships using noble metal anodes and the ship's own electrical system.
  A reference cell such as silver/silver chloride is mounted to detect voltage differences between itself and the hull.
  This voltage difference is measured and a reactor rectifier is used to control the current from the ship's electrical system to the external anodes in order to prevent corrosion.

• Alloying with Another Metal

  Alloying iron with chromium produces stainless steel.
  The chromium readily forms an impermeable (impervious) oxide layer that adheres strongly to the iron surface preventing the formation of rust.