The Science of Corrosion

The Science of Corrosion

Introduction

corrosion-on-ships1We live and work with metals, we have done for thousands of years, we see how some metals change colour and almost disintegrate before our eyes. The nature of corrosion of metal comes from the metal’s interaction with its environment with specific environmental and physical conditions, and the process is a one way street to destruction if left to its own devices. Economically, the cost of damage due to corrosion, directly and indirectly, is enormous, estimates of around 4% of the world’s GDP are frequently quoted and many statistical studies have pretty much been in agreement. Consequently, the market for corrosion control measures is equally as large and comprise of many differing methods to act as a barrier to the environment.

The Study of Corrosion

The study of corrosion is actually a pairing of two disciplines: Corrosion science and Corrosion engineering. Corrosion science seeks to understand the underlining mechanism of corrosion, through scientific knowledge in the areas of chemistry and material science, including mechanics. Corrosion engineering takes that knowledge gained through scientific endeavour and applies that in a practical way, in order to provide the necessary protection. This might give the impression that the two fields of study are separate and independent of each other, experience has shown that this true to an extent, however it need not be and it is preferable that the two work hand-in-hand together.

The relationship between the sciences can best be summed up as:

Environment + Metal + Conditions = Type of Corrosion under investigation

The Environment

Under the environment tag, the number of different environmental conditions could include:

  • The ocean
  • Fresh Water lakes and rivers
  • The atmosphere
    • Industrial
    • Marine
    • Indoor
    • Cities
  • The earth or soils
  • Storage tanks or vessels
  • Pipelines

The Metal

Under the metal tag, the common variable is the chemistry of the alloy; the effect of different metal compositions can have a profound affect on the rate of corrosion. For example the addition of Chromium and Nickel into steel creates a very tough oxide surface layer, which acts as a barrier between the Iron in steel and the environment. Contrast that to Aluminium, for example, which naturally forms a tough oxide surface layer in contact the atmosphere.

The Conditions

Under the conditions tag, the physical conditions in which the metal is been subjected to. This not only includes the chemical conditions, like concentrations, presence of oxygen and other gases and the presence of chemical substances, but it also includes mechanical and physical conditions like temperature, pressure, stress and fluid dynamics.

 

What is Corrosion

rusting-eyeletSimply put, corrosion is a destructive chemical reaction of the metal with its environment. We commonly think of corrosion as simply rust, that flaky red deposit on a steel surface, and while that is a form of corrosion, it is just one example of such. In our case, the descriptive word “corrosion” is expanded to refer to the destructive chemical reaction of any metal by its environment. Rust applies to the corrosion of iron and mild steels by the presence of oxygen in water in contact with the metal surface. The product of this chemical reaction, called rust, is a hydrated iron oxide and makes its appearance as red deposits.

To be much more exact, corrosion is an electrochemical reaction. That is, corrosion is a result of the pairing or coupling of electrochemical half-cell reactions – A REDOX reaction or a Reduction Oxidation reaction.

Half-cell reactions involves the transfer of electrons from one side of the reaction to the other. If the electrons transfer to the products, indicating a loss of electrons, then it is an oxidation reaction. For example:

Fe(s) → Fe2+(aq) + 2e

 

If the electrons form part of the reactants, then it is a reduction reaction. for example:

2H+(aq) + 2e− → H2(g)

So combining the two equations to form an electrochemical, we can see that the electrons from the oxidation reaction reduce the H+ (typically acidic solutions) to hydrogen gas:

Fe(s) + 2H+(aq) → Fe2+(aq) + H2(g)

On a metal surface, these half-cell reactions can take place at a number of different sites, with the electrons being conducted elsewhere. The reason why these reactions take place, the reason why there is a potential difference allowing electrons to flow, lies with the structure of the metal. Metal may look smooth and continuous, but the crystalline structure features an array of defects, grain boundaries, dislocations and impurities that leads to localised potential differences. If the metal is in contact with an ionic solution, sea water is a good example with dissolved sodium and chloride ions, the ions absorb into the surface of the metal and that can change electrochemical reaction. It is this ionic solution that can act as an electrolyte to conduct a current or flow of electrons from one site to another.

Seawater typically contains about 3-4% dissolved sodium chloride, which unfortunately corresponds to a maxima in the rate of corrosion. In comparison to fresh water, sea water is far more corrosive to iron and mild steels, due to its conductivity but also to the penetrating power of the chloride ion.

Therefore, for corrosion to take place four conditions are necessary, and these are:

  1. An oxidation reaction
  2. A reduction reaction
  3. A metallic conduction path between the the half-cell reactions
  4. The presence of an electrolyte

Looking back at the half-cell reactions for iron, we can see that iron is oxidised to iron II (Fe2+), but to create rust (Iron Oxide) we need to reduce the dissolved oxygen in water to hydroxide ions (OH-):

O2 (g) + 2 H2O (l) + 4e–  → 4 OH–

The resulting reaction produces Iron Hydroxide, which is futher oxidised in the presence of dissolved oxygen in water to Iron (III) Oxide – Rust!

As a note of interest, note that hydroxide ions are produced from dissolved oxygen in water and accumlate as the reaction progresses. This causes the pH of the thin layer on the surface of the metal to rise (pH>7), i.e. the solution becomes more alkaline and with changes in pH the rate of corrosion changes dramatically. Aluminium, for example, has a very low reaction rate at around pH7, but increases by a factor of X 1,000 at pH 14, this is evident when sodium hydroxide (Causti Soda) comes in contact with aluminium foil or kitchen pans commonly found in kitchens.