Corrosion and Prevention: Definition of Corrosion, Rusting, Mechanism of Corrosion, Factors Assisting Corrosion, Prevention of Corrosion, Practice Problems and FAQs:

Have you ever been to the technology marvel of India- the Qutub minar? 

Well, at least you must have seen videos and photos of it. 

What is the most exciting part of this famous minar? It is an iron (mostly of iron) pillar standing in the hot Sun and thunder storms but without any rusting for more than 300 years. 

Take a closer look at the picture of the iron pillar, now look at the bicycle which you bought last year. Don’t be sad about it. 

Your curious brain must be wondering how your bicycle caught rust even when you are keeping it away from rain and heat? While the pillar which is left open to air, toxic pollution (obviously it is situated in Delhi), moisture everything that can cause rusting is still looking new?

If you can find the secret behind it, you will not only save billions worth mineral being wasted but earning billions of dollars. A challenge for the younger generation!

Well, let’s understand why rusting happens, what is the relation between rusting and corrosion and what are the methods to protect metals from rusting. 

Table of content

• Definition of corrosion
• Mechanism of corrosion.
• Factors which promote corrosion :
• Prevention of corrosion
• Practice problems
• Frequently asked questions (FAQs)

Definition of corrosion

Corrosion is the natural destruction of metals due to their interaction with the environment. On the exposed surface, corrosion occurs. When the metal's upper layer is corroded or removed, the inner surface of the metal gets exposed, and the corrosion continues to a certain depth.

In other words, the process of slowly eating away of the metal due to attack of the atmospheric vapours on the surface of the metal resulting into the formation of compounds such as oxides, sulphides, carbonates, sulphates etc. is called corrosion.

• As soon as the metals are extracted from their ores, the reverse process begins, i.e., nature tries to convert them back into the form in which they occur. This is due to the attack of the gasses present in the atmosphere on the surface of the metal.

• Some examples include tarnishing of silver, development of green coating on copper and bronze, white layer on aluminum, etc. 
• Rusting- Rusting is related to corrosion on iron metal. The most common example of corrosion is the rusting of iron. Rust is hydrated ferric oxide, Fe₂O₃.xH₂O

Mechanism of corrosion

The theory of corrosion can be explained by taking the example of rusting of iron. The theory is called electrochemical theory because it explains the formation of rust on the basis of the formation of electrochemical cells on the surface of the metal.

The formation of rust on the basis of this theory may be explained in the following steps:

• The water vapours on the surface of the metal dissolve CO₂ and O₂ from the air. Thus, the surface of metal is covered with the solution of CO₂ in water, i.e., carbonic acid (H_₂CO₃)

H₂O + CO₂ → H₂CO₃……………….(i)

This acts as an electrolytic solution of the cell. The week carbonic acid dissociates to a small extent as follows:

H₂CO₃⇌ 2H⁺+CO₃^2-……………….(ii)

H+ ions may also be produced by dissolution of other acidic oxides such as SO₂, Cl₂, liberated from industries, present in the atmosphere.

• Iron in contact with the dissolved Oxygen, and in presence of H+ ions undergoes oxidation as follows: 

FeFe²⁺+2^e- E⁰_oxidation=0.44 V………(iii)

Thus, the sites where the above reaction takes place act as anodes. As a result of the above reaction, iron is converted into ferrous (Fe²⁺) ions.

• The electrons lost by iron migrate through the metal to another site on the surface and reduce dissolved oxygen in the presence of H⁺ ions as follows:

O_₂ + 4^H+ + 4 ^e- 2H^_₂O (E0reduction= 1.23 V)…………..(iv)

The sites where the above reaction takes place act as cathodes.

Notice that H+ ions are involved in the reduction of O₂. As the concentration of H⁺ ions is lowered (i.e., pH is increased), the reduction of oxygen becomes less favourable. It is observed that iron in contact with a solution whose pH is above 9-10 does not corrode.

Adding equations (iii) and (iv), the overall reaction of the miniature cell will be

2 Fe (s) + O₂(g) + 4^H+ (aq) →2Fe²⁺ (aq) + 2H₂O (l); E⁰cell = 1.67 V

It may be mentioned here that if water is saline, it increases the acidity of water around that helps in the flow of current in the miniature cell and hence enhancing the process of corrosion. It is for this reason that in places where it snows heavily in winter and salting of roads is done, the automobile vehicles on the road undergo greater rusting.

• The ferrous ions formed react further with the dissolved oxygen or oxygen from the air to form ferric oxide with further production of H+ ions as follows:

4Fe²⁺+O₂+4H₂O2Fe₂O₃+8H⁺(aq)

•  Ferric oxide then undergoes hydration to form rust as follows:

Fe₂O₃+xH₂OFe₂O₃.xH₂O (Hydrated ferric oxide or rust)

It is interesting to note that as cathode is generally the area having the largest amount of dissolved oxygen, the rust is often formed here. If you look at any rusted article, you will observe that rust is

formed at a site other than the site where pitting (Pitting is a type of corrosion that happens in materials with protective films. It is a surface-based attack with localized holes. The attack can quickly penetrate the metal, while some parts of the metal surface remain corrosion-free) has occurred.

It may be noted that the rust is a non-sticking compound, i.e., it does not stick to the surface. It peels off

exposing fresh iron surface for further rusting. 

Factors which promote corrosion :

• Reactivity of the metal. More active metals are readily corroded.
• Presence of impurities. Presence of impurities in metals enhances the chances of corrosion. Pure metals do not corrode, e.g., pure iron does not rust.
• Presence of air and moisture. Air and moisture accelerate corrosion. Presence of acidic gasses like SO2, and CO, in air catalyse the process of corrosion. Iron when placed in vacuum does not rust.
• Strains in metals. Corrosion (e.g., rusting of iron) takes place rapidly at bends, scratches, nicks and cuts in the metal.
• Presence of electrolytes. Electrolytes, if present, also increase the rate of corrosion. For example, iron rusts faster in saline water than in pure water.

Prevention of corrosion

Prevention of corrosion is very important because enormous damage occurs to buildings, bridges, ships and all other objects made of metals especially iron. Crores of rupees are spent every year on replacing the rusted structures. Prevention of corrosion not only saves money but also helps in preventing accidents such as bridge collapse etc.

Some of the methods generally used for prevention of corrosion are briefly explained below :

• Barrier Protection

The metal surface is not allowed to come in contact with moisture, oxygen and carbon dioxide. This can be achieved by the following methods:

• Painting: 

The metal surface is coated with paint which keeps it out of contact with air, moisture etc. till the paint layer develops cracks.

• Oil and grease:

By applying film of oil and grease on the surface of the iron tools and machinery, the rusting of iron can be prevented since it keeps the iron surface out of contact with moisture, oxygen and carbon dioxide.

• Metal coating:

The iron surface is coated with non-corroding metals such as 

• nickel, chromium, aluminum, etc. by electroplating,

•  zinc( called galvanisation.) by dipping the iron article in the molten active metal. The layer of zinc on the surface of iron, for example when comes in contact with moisture, oxygen and carbon-dioxide in air, forms a protective invisible thin layer of basic zinc carbonate. ZnCO.Zn(OH) iron s lose its lustre and gets protected from further corrosion.

This again shuts out the supply of oxygen and water to iron surface.

• Iron can be coated with copper by electrodeposition from a solution of copper sulphate or with tin by dipping into molten metal. Now, if the coating is broken, iron is exposed and iron being more active than both copper and tin, is corroded. Here, iron corrodes more rapidly than it does in the absence of tin At times, zinc, magnesium and aluminum powders mixed with paints provide decorative protective coatings also

• Using anti-rust solutions

The iron surface is coated with phosphate or other chemicals which give a tough adherent insoluble film which does not allow air and moisture to come in contact with iron surface.

These are alkaline phosphate and alkaline chromate solutions. The alkalinity prevents the availability of hydrogen ions. In addition, phosphate tends to deposit an insoluble protective film of iron phosphate on the iron. These solutions are used in car radiators to prevent rusting of iron parts of the engine.

• Alloy formation-

Alloys have improved properties than the basic metal. Iron for example is alloyed with chromium, carbon, etc t prepare variety of steels that are more resistant to corrosion. Stainless steel we use as utensils is nothing but a low carbon iron alloy.

• Passivation-

Some metal surfaces on treatment with oxidizing agents gets oxidized to form oxides. This surface layer of metal oxide does not undergo further chemical reaction and by that protect the bottom layers from corrosion. Tis is called passivation of the metal surface.

Such passivation occurs with atmospheric oxygen but rather slowly. Aluminum metal for example forms a white surface layer of aluminum oxide that prevents further corrosion of aluminum metal

• Sacrificial Anode-Cathode Protection

Sacrificial protection refers to the loosing (sacrificing) an another metal to protect a perticular metal of interest. This is done by coupling the metal to be protected to an another more electropositive metal sp a sto form a electrical circuit.. Higher the electropositive nature of the metal higher the tendency of the metal to loose electrons and get oxidized than a lower electropositive metal.

With the passage of time, the more active metal gets consumed but so long as it is present there, it will protect the metal from corrosion and does not allow even the nearly exposed surface to lose electrons (i.e., undergo oxidation). 

The iron object acts as cathode and the protecting metal acts as anode. The anode is gradually used up due to the oxidation of the metal to its ions due to loss of electrons. These electrons are transferred through to H+ ions present around the iron object and thus protect it from rusting. As iron object acts as cathode, it is called cathodic protection. The iron object gets protection from rusting as long as some of the active metal is present. Metals widely used for protecting iron objects from rusting are magnesium, zinc and aluminum which are sacrificial anodes.

Magnesium is often employed in the cathodic protection of iron pipes buried in the moist soil, canals, storage tanks, etc. Pieces of magnesium are buried along the pipeline and connected to it by the wire.

Practice problems 

Q1. Iron gets corroded to form rust. Rust contains\

a. FeO and Fe(OH)₂
b. FeO and Fe(OH)₃
c. Fe₂O₃ and Fe(OH)₃
d. Fe₃O₄ and Fe(OH)₃

Answer: (C)
Rust is generally a mixture of Fe₂O₃ and Fe(OH)₃.

Q2. In galvanization, coating of zinc on iron is done to protect iron. But iron coated on zinc gets rusted because- 

a. zinc is lighter than iron
b. zinc has lower melting point than iron
c. zinc has lower reduction electrode potential than iron
d. zinc has higher negative reduction potential than iron

Answer: (D)
Zinc has higher negative electrode potential than iron, so iron can’t be coated on zinc

Q3. Which of the given processes will increase instead of prevention of corrosion?

a. greasing
b. painting
c. Heating
d. plating

Answer: (C)
Painting, oiling, greasing, or varnishing the surface of iron will protect it from rusting. Another way of keeping iron from rust is galvanisation, which involves coating it with a thin layer of zinc. Iron corrosion is prevented by coating it with a non-corrosive substance such as carbon.

Q4. Choose the correct statement on Corrosion?

a. Corrosion is a natural process that converts a metal into a more stabler form
b. It is the gradual destruction of materials (usually a metal) by reaction with their environment
c. Corrosion science is the study of controlling and preventing corrosion
d. All of the above

Answer: (D)
Corrosion is a natural process that transforms a refined metal into a more chemically stable form like oxide, hydroxide, or sulphide. It refers to the progressive decomposition of things (typically metals) due to chemical and/or electrochemical reactions with their surroundings. Corrosion engineering is the science of preventing and controlling corrosion.

Frequently asked questions (FAQs)

Q1. Does rusting requires air?
Answer: Yes. Rusting is a form of oxidation. Rust is formed when iron reacts with water and oxygen to form hydrated iron(III) oxide. When iron and steel come into contact with water and oxygen (both of which are required for rusting to occur), they rust.

Q2. The rust formed during the process of rusting is soluble in water or not?
Answer: Insoluble. This is unusual because iron oxides and iron hydroxides (components of rust) do not normally dissolve in water, so the water is reduced while the iron is oxidized.

Q3. Why does the color of rust appear to be red?
Answer: Uniform corrosion is the most common type of corrosion. This is when the material's surface is covered in an even layer of oxidation. When metal is exposed to a large amount of water and oxygen, the iron reacts with a contaminate, resulting in "red" rust.