Corrosion

Corrosion is one of the most common phenomena that we observe in our daily lives.  You must have noticed that some objects made of iron are covered with some orange or reddish-brown coloured layer at some point in time. The formation of this layer is the result of a chemical process known as rusting, which is a form of corrosion.

Corrosion can be defined as the process through which refined metals are converted into more stable compounds such as metal oxides, metal sulfides, or metal hydroxides. The rusting of iron involves the formation of iron oxides via the action of atmospheric moisture and oxygen. Corrosion is usually an undesirable phenomenon since it negatively affects the desirable properties of the metal. For example, iron is known to have good tensile strength and rigidity (especially alloyed with a few other elements). However, when subjected to rusting, iron objects become brittle, flaky, and structurally unsound. Corrosion can be classified as an electrochemical process since it usually involves redox reactions between the metal and certain atmospheric agents such as water, oxygen, sulphur dioxide, etc.

If we look at the science behind corrosion then we can say that it is a spontaneous/ irreversible process wherein the metals turn into a much stable chemical compound like oxides, sulphides, hydroxides, etc. We shall delve deeper into the concept of corrosion and understand its different factors including its meaning, types, prevention and more in this lesson.

Table of Content

Corrosion Definition

What is Corrosion? It is defined as the natural process that causes the transformation of pure metals to undesirable substances when they react with substances like water or air. This reaction causes damage and disintegration of the metal starting from the portion of the metal exposed to the environment and spreading to the entire bulk of the metal.

Do All Metals Corrode?

Metals placed higher in the reactivity series such as iron, zinc, etc. get corroded very easily and metals placed lower in the reactivity series like gold, platinum and palladium do not corrode. The explanation lies in the fact that corrosion involves oxidation of the metals. As we go down the reactivity series tendency to get oxidised is very low (oxidation potentials is very low).

Also read: Oxidation and Reduction

Interestingly, aluminium doesn’t corrode unlike other metals even though it is reactive. This is because aluminium is covered by a layer of aluminium oxide already. This layer of aluminium oxide protects it from further corrosion.

Factors Affecting Corrosion

1. Exposure of the metals to air containing gases like CO2, SO2, SO3 etc.

2. Exposure of metals to moisture especially salt water (which increases the rate of corrosion).

3. Presence of impurities like salt (eg. NaCl).

4. Temperature: An increase in temperature increases corrosion.

5. Nature of the first layer of oxide formed: some oxides like Al2O3 forms an insoluble protecting layer which can prevent further corrosion. Others like rust easily crumble and expose the rest of the metal.

6. Presence of acid in the atmosphere: acids can easily accelerate the process of corrosion.

Types of Corrosion

Some of the corrosion types include;

(i) Crevice Corrosion

Whenever there is a difference in ionic concentration between any two local areas of a metal, a localized form of corrosion know as crevice corrosion can occur. Examples of areas where crevice corrosion can occur are gaskets, the undersurface of washers, and bolt heads.

Example: All grades of aluminium alloys and stainless steels undergo crevice corrosion.

(ii) Stress Corrosion Cracking

Stress Corrosion Cracking can be abbreviated to ‘SCC’ and refers to the cracking of the metal as a result of the corrosive environment and the tensile tress placed on the metal. It often occurs at high temperatures.

Example: Stress corrosion cracking of austenitic stainless steel in chloride solution.

(iii) Intergranular Corrosion

Intergranular corrosion occurs due to the presence of impurities in the grain boundaries that separate the grain formed during the solidification of the metal alloy. It can also occur via the depletion or enrichment of the alloy at these grain boundaries.

Example: Aluminum-base alloys are affected by IGC.

(iv) Galvanic Corrosion

When there exists an electric contact between two metals that are electrochemically dissimilar and are in an electrolytic environment, galvanic corrosion can arise. It refers to the degradation of one of these metals at a joint or at a junction. A good example of this type of corrosion would be the degradation that occurs when copper, in a salt-water environment, comes in contact with steel.

Example: When aluminium and carbon steel are connected and immersed in seawateraluminium corrodes faster and steel is protected.

(iv) Pitting Corrosion

Pitting Corrosion is very unpredictable and therefore is difficult to detect. It is considered one of the most dangerous types of corrosion. It occurs at a local point and proceeds with the formation of a corrosion cell surrounded by the normal metallic surface. Once this ‘Pit’ is formed, it continues to grow and can take various shapes. The pit slowly penetrates metal from the surface in a vertical direction, eventually leading to structural failure it left unchecked.

Example: Consider droplet of water on steel surface, pitting will initiate at the centre of the water droplet (anodic site).

(v) Uniform Corrosion

This is considered the most common form of corrosion wherein an attack on the surface of the metal is executed by the atmosphere. The extent of the corrosion is easily discernible. This type of corrosion has a relatively low impact on the performance of the material.

Example:  A piece of zinc and steel immersed in diluted sulphuric acid would usually dissolve over its entire surface at a constant rate.

Corrosion Examples and Reactions

Here are some of the typical examples of corrosion as seen mostly in metals.

1. Copper Corrosion

When copper metal is exposed to the environment it reacts with the oxygen in the atmosphere to form copper (I) oxide which is red in colour.

2Cu(s) + ½ O2(g) → Cu2O(s)

Cu2O further gets oxidised to form CuO which is black in colour.

Cu2O(s) + ½ O2(g) → 2CuO(s)

This CuO reacts with CO2, SO3 and H2O (present in the atmosphere to form Cu2(OH)2(s) (Malachite) which is blue in colour and Cu4SO4(OH)6(s) (Brochantite) which is green in colour.

This is why we observe copper turning bluish-green in colour.

A typical example of this is the colour of the statue of liberty which has the copper coating on it turning blue-green in colour.

2. Silver Tarnishing

Silver reacts with sulphur and sulphur compounds in the air give silver sulphide (Ag2S) which is black in colour. Exposed silver forms Ag2S as it reacts with the H2S(g) in the atmosphere which is present due to certain industrial process.

2Ag(s) + H2S(g) → Ag2S(s) + H+2+(g)

3. Corrosion of Iron (Rusting)

Rusting of iron which is the most commonly seen example happens when iron comes in contact with air or water. The reaction could be seen as a typical electrochemical cell reaction. Consider the diagram given below.

Corrosion of Iron (Rusting)

Here metal iron loses electrons and gets converted to Fe{aq}2+ (this could be considered as the anode position). The electrons lost will move to the other side where they combine with H+ ions. H+ ions are released either by H2O or by H2CO3 present in the atmosphere(this could be considered as the cathode position).

H2O \rightleftharpoons H+ + OH

H2CO3 \rightleftharpoons 2H+ + CO32

The Hydrogen thus formed by the reaction of H+ and electrons react with oxygen to form H2O.

Also Read: Rusting of Iron and Prevention

Anode reaction

2Fe(s) → 2Fe2+ + 4eεFe2+/Fe0\varepsilon _{_{F{{e}^{2+}}/Fe}}^{0} = – 0.44 V

Cathode reaction

O2(g)+4H+(aq)+4e2H2O(l)EoH+/O2/H2/O=1.23V{{O}_{2(g)\,}}+4{{H}^{+}}_{(aq)}+4{{e}^{-}}\overset {2} \longrightarrow {{H}_{^{2}}}{{O}_{(l)}}{{E}^{o}}_{^{{{H}^{+}}}/{{O}_{2}}/{{H}_{2}}/O\,\,\,}=1.23V

Overall reaction

2Fe(s) + O2(g) + 4H+(aq) → 2Fe2+(aq) + 2H2O(l)Eocell = 1.67V

The Fe2+ ions formed at the anode react with oxygen in the atmosphere, thereby getting oxidised to Fe3+ thereby forming Fe2O3 which comes out in the hydrated form as Fe2O3.xH2O

Fe2+ + 3O2 → 2Fe2O3

Fe2O3 + xH2O → Fe2O3. xH2O (rust)

Other examples include,

  • Corrosion of Zinc when it reacts with oxygen and HCl to form white coloured ZnCl2.
  • Corrosion of Tin to form black coloured Na2[Sn(OH)2].

Prevention of Corrosion

Preventing corrosion is of utmost importance in order to avoid huge losses. Majority of the structures we use are made out of metals. This includes bridges, automobiles, machinery, household goods like window grill, doors, railway lines, etc. Some of the popular methods to prevent corrosion include Electroplating, Galvanization, Painting and Greasing, Use of Corrosion Inhibitor, etc.

Read More: How Can Corrosion Be Prevented

BOOK

Free Class