Wet (or) Electro-chemical Corrosion

WET OR ELECTRO-CHEMICAL CORROSION

Wet corrosion occurs under the following conditions.

(i) When two dissimilar metals or alloys are in contact with each other in the presence of an aqueous solution or moisture.

(ii) When a metal is exposed to varying concentration of oxygen or any electrolyte.

Mechanism of wet corrosion

Under the above conditions, one part of the metal becomes anode and another part becomes cathode.

1. At anode

In anodic part, oxidation (or) dissolution of metal occurs

M → M²⁺ + 2e^‒

2. At cathode

In cathodic part reduction reaction occurs, which depends on nature of the corrosive environment.

(a) Acidic environment

If the corrosive environment is acidic, hydrogen evolution occurs at cathodic part.

2H⁺ + 2e^‒→ H₂↑

(b) Neutral environment

If the corrosive environment is slightly alkaline (or) neutral, hydroxide ions are formed at cathodic part.

½ O₂ + 2e^‒ + H₂O  → 2OH^‒

Thus, the metal ions (from anodic part) and non‒metallic ions (from cathodic part) diffuse towards each other through conducting medium and form a Corrosion product between anode and cathode.

(a) Hydrogen evolution type corrosion

"All metals above hydrogen in the electrochemical series have a tendency to get dissolved in acidic solution with simultaneous evolution of hydrogen gas" (Fig. 4.2).

Example: When iron metal contacts with non‒oxidising acid like HCl, H₂ evolution occurs.

At anode

Iron undergoes dissolution to give Fe²⁺ ions with the liberation of electrons.

Fe → Fe²⁺ + 2e^‒ (oxidation)

At cathode

The liberated electrons flow from anodic to cathodic part, where H⁺ ions get reduced to H₂.

2H⁺ + 2e ̄  → H₂↑ (reduction)

(b) Absorption of oxygen (or) Formation of hydroxide ion type corrosion

The surface of iron is usually, coated with a thin film of iron oxide. However, if the oxide film grows, some crack will form and anodic areas are created on the surface while the remaining part acts as cathode (Fig. 4.3).

Example:

When iron metal contacts with a neutral (or) slightly alkaline solution of an electrolyte in presence of oxygen, OH^‒ ions are formed.

At anode

Iron, dissolves as Fe²⁺ with the liberation of electrons.

 Fe →  Fe²⁺ + 2e^‒ (oxidation)

At cathode

The liberated electrons flow from anodic to cathodic part through metal, where the electrons are taken up by the dissolved oxygen to form OH^‒ ions.

 ½O₂ + H₂O + 2e^‒ → 2OH^‒

Thus, the net corrosion reaction is

 Fe²⁺ + 2OH ̄  → Fe(OH)₂↓

If enough O₂ is present, Fe(OH)₂ is easily oxidized to Fe(OH)₃, a rust (Fe₂O₃. H₂O).

4Fe(OH)₂ + O₂ + 2H₂O → 4Fe(OH)₃ (or)

(2Fe₂O₃.3H₂O) brown rust

Table 4.1 Difference between Chemical and Electro chemical Corrosion

Chemical Corrosion

1. It occurs only in dry condition.

2. It is due to the direct chemical attack of metal by the environment.

3. Even a homogeneous metal surface gets corroded.

4. Corrosion products accumulate in the same place, where corrosion occurs.

5. Chemical corrosion is self‒controlled.

6. It follows adsorption mechanism.

7. Examples: Formation of mild scale on iron surface.

Electrochemical Corrosion

1. It occurs in presence of moisture or electrolyte.

2. It is due to the set up of a large number of cathodic and anodic areas.

3. Heterogeneous surface (or) Bimetallic contact is the condition.

4. Corrosion occurs at the anode, while products formed elsewhere.

5. Electrochemical corrosion is continuous process.

6. It follows electrochemical reaction.

7. Examples: Rusting of iron in moist atmosphere.

1. Types of electrochemical corrosion

1. Galvanic corrosion

When two different metals are in contact with each other in presence of an aqueous solution or moisture, galvanic corrosion occurs.

Here, the more active metal (with more negative electrode potential) acts as anode and the less active metal (with less negative electrode potential) acts as cathode.

Example: Zn‒Fe couple

Fig. 4.4(a) represents Zn‒Fe couple, in which zinc (more active or higher in emf series) dissolves in preference to iron (less active metal) i.e., Zn acts as anode and undergoes corrosion and Fe acts as cathode.

Example: Cu‒Fe couple

Fig. 4.4 (b) represents Fe ‒ Cu couple, in which iron (more active, when compared to Cu) dissolves in preference to copper (less active) i.e., Fe acts as anode and undergoes corrosion and Cu acts as cathode.

Examples for Galvanic Corrosion

(i) Steel screw in a brass marine hardware corrodes

This is due to galvanic corrosion. Iron (higer position in electrochemical series) becomes anodic and is attacked and corroded, while brass (lower in electrochemical series) acts as cathodic and is not attacked.

(ii) Bold and Nut made of the same metal is preferred

It is preferred in practice, because galvanic corrosion is avoided due to homogeneous metals (no anodic and cathodic part).

Prevention

Galvanic corrosion can be minimised by the following two ways.

1. By providing an insulating material between the two metals.

2. By selecting two metals as close as possible on the emf series.

2. Differential aeration (or) concentration cell corrosion

This type of corrosion occurs when a metal is exposed to varying concentration of oxygen or any electrolyte on the surface of the base metal.

Example: Metals partially immersed in water (or) conducting solution (called water line corrosion).

If a metal is partially immersed in a conducting solution (Fig. 4.5) the metal part above the solution is more aerated and hence become cathodic. On the other hand, the metal part inside the solution is less aerated and thus, become anodic and suffers correcsion.

At anode (less aerated part): Corrosion occurs

M → M²⁺ + 2e^‒

At cathode (more aerated part): OH^‒ ions are produced

 ½O₂ + H₂O + 2e^‒ → 2OH^‒

Examples for Differential aeration corrosion

(a) Pitting or localised corrosion.

(b) Crevice corrosion.

(c) Pipeline corrosion.

(d) Corrosion on wire fence.

(a) Pitting corrosion

Pitting is a localised attack, resulting in the formation of a hole around which the metal is relatively unattacked.

Example: Metal area covered by a drop of water, sand, dust, scale, etc.

Let us consider a drop of water or aqueous NaCl resting on a metal surface (Fig. 4.6). The area covered by the drop of water acts as an anode due to less oxygen concentration and suffers corrosion. The uncovered area (freely exposed to air) acts as a cathode due to high oxygen concentration.

The rate of corrosion will be more, when the area of cathode is larger and the area of anode is smaller, Therefore, more and more material is removed from the same spot. Thus a small hole or pit is formed on the surface of the metal.

At anode

Iron is oxidised to Fe²⁺ ions

Fe → Fe²⁺ + 2e^‒

At cathode

Oxygen is converted to OH^‒ ions.

½O₂ + H₂O + 2e ̄ → 2OH^‒

Net reaction

Fe²⁺ + 2OH^‒ → Fe(OH)₂ –^(O)→ Fe(OH)₃

This type of intense corrosion is called pitting.

(b) Crevice corrosion

If a crevice between different metallic objects of between metal and non‒metallic material is in contact with liquids, the crevice becomes the anodic region and suffers corrosion. This is due to less oxygen in crevice area. The exposed areas act as cathode (Fig. 4.7).

(c) Pipeline corrosion

Differential aeration corrosion may also occur in different parts of pipeline.

Buried pipelines or cables passing from one type of soil to another say, from clay (less aerated) to sand (more aerated) may get corroded due to differential aeration.

(d) Corrosion on wire‒fence

Fig. 4.9 shows a wire fence in which the areas where the wires cross are less aerated than the rest of the fence and hence corrosion occurs at the wire Crossings, which are anodic.

Other examples for differential aeration corrosion

1. Corrosion occurring under metal washers, where oxygen cannot diffuse easily.

2. Lead pipeline passing through clay to cinders undergo corrosion. Since the pipeline under cinders is more aerated, it gets corroded easily.