Corrosion of Materials

Engineering Materials 2189101

Chedtha Puncreobutr

Department of Metallurgical Engineering Chulalongkorn University

Corrosion of Materials

Things we've learnt about oxidation

• In dry condition, most

materials that are unstable in oxygen tend to oxidized

• Oxidation is controlled by diffusion of ions or the conduction of electrons through oxide films

• Diffusion is thermally activated process thus rate of oxidation is much greater at high

temperature

• Loss of materials by oxidation

at room temperature under

dry conditions is very slight

Durability of materials

Degradation: Reduction in performance of component during lifetime

Mechanical processes

Chemical processes

Creep

Oxidation

Electrochemical Corrosion Fatigue

Wear

Electrochemical (“wet”) Corrosion

• It can simply occur due to water in atmosphere

• Also it can be promoted by

• acids, alkalis, and salts

• sea water

• presence of other

Comparing to oxidation, the only real difference is the medium that causes the process. Oxygen causes oxidation, while corrosion is the term applied to a similar electrochemical process caused by many

Corrosion under wet conditions

Electrochemical (“wet”) Corrosion

Economic impact: ~3.5% of GDP in developed countries

Direct losses

Component replacement

Painting and other preventative measures

Use of expensive corrosion resistant materials

Indirect losses

Plant shutdown

Loss of product e.g. in pipes Loss of efficiency

Contamination of product

Without prevention, corrosion can cause catastrophic 1992 Guadalajara explosions

• Underground humidity caused these materials to create an electrolytic reaction, Steel gas pipe then corroded, creating a hole in the pipeline that permitted gas to leak into the ground and into the main sewer pipe.

• 252 people were killed

• 500 injured

• 15,000 were left homeless.

• Monetary damage

ranges between $300

million and $1 billion.

Wet Corrosion

Materials in aqueous solutions tend to form ions

𝑀 → 𝑀 ++ + 2𝑒 −

Fe Fe⁺⁺

2e^-

Fe Fe⁺⁺

2e^-

Abraded Iron Aerated water

Pass into the water Leaving behind electrons

Wet Corrosion

Fe Fe⁺⁺

2e^-

Fe Fe⁺⁺

2e^-

Abraded Iron Aerated water

The electrons are then conducted through the metal to a place where oxygen reduction reaction can take place to consume electrons (cathodic reaction)

𝑂₂ + 2𝐻₂𝑂 + 4𝑒⁻ → 4𝑂𝐻⁻ Cathode

(Reduction)

Anode (Oxidation)

This reaction generates OH

^-

ions

Wet Corrosion

𝑂₂ + 2𝐻₂𝑂 + 4𝑒⁻ → 4𝑂𝐻⁻

Fe Fe⁺⁺

Fe Fe⁺⁺

Abraded Iron Aerated water 𝐹𝑒 𝑂𝐻 ₂ 𝑖. 𝑒. 𝐹𝑒𝑂 ∙ 𝐻₂𝑂

conductive solution (electrolyte)

The OH

^-

ions then combine with Fe

⁺⁺

ions to form a hydrated iron oxide Fe(OH)

₂

Cathode (Reduction)

Anode (Oxidation)

Material + Oxygen  (Hydrated) Material Oxide

Wet Corrosion

• The Fe(OH)

₂

either deposits away or loosely deposits on the surface, giving little or no protection

• Very rapid diffusion as M

⁺⁺

and OH

^-

usually diffuse in liquid

• In conducting materials, electrons can move very easily as well Much faster attack

found in wet corrosion

compared to dry air

Voltage differences as the driving force

The tendency of a metal to oxidized in solution (wet corrosion) is described by using a voltage scale rather than energy

• Oxygen is reduced

• Absorbing electrons

• Metal becomes positively charged

• Reaction continues until potential raise to +0.8V

• Fe++ forms

• Leaving electrons behind

• Acquires negatively charged

• Reaction continues until potential fall to -0.6V

Voltage differences as the driving force

The tendency of a metal to oxidized in solution (wet corrosion) is described by using a voltage scale rather than energy

• If the anode and

cathode are connected, electron flow from one to the other.

• The difference in voltage of 1.4V is driving force for oxidation reaction

Bigger voltage difference

leads to bigger tendency

to corrode

Pourbaix (Electrochemical equilibrium) Diagram

The diagrams are maps that show the conditions under which a metal:

• Cannot corrode (immunity)—

because there is no voltage driving force, or a negative one

• May corrode (corrosion)—

because there is a voltage driving force, and a stable oxide film does not form on the surface

• May not corrode (passivation)—

although there is a voltage

driving force, a stable oxide film forms on the surface (this may or may not be an effective barrier to corrosion)

Corrosion in Copper

• A protective film has

formed which is why the colour is darker than freshly polished copper

• Bright green colour in old architecture because the carbon dioxide in the air forms a copper carbonate

• Copper is used as water pipe because corrosion rate is extremely low when water is neutral or alkaline (oxide film acted as an effective barrier)

• If the water is mildly acidic, corrosion “pinholes” can form in the tube wall

Corrosion in Copper alloys

• Brass (copper and zinc) and Bronze (copper and tin) alloys are resistant to corrosion

Brass fire hydrant

Corrosion in Steel

• To have stable surface film, the pH of the water needs to be above 9

• so in most applications (neutral or slightly acidic water) steel will rust unless it is protected.

• Making steel resistant to corrosion is to alloy it with foreign elements

• Stainless steel - adding 18% chromium produces an invisible film of Cr₂O₃

• Another way is to keep any moisture and/or oxygen away from the surface of the steel (e.g., using paint or epoxy coatings).

Corrosion in Aluminium

• Between pH 4 and 8.5, a thin

and very stable film of hydrated aluminium oxide forms, protecting the metal

• The corrosion rate of

aluminium in pure water is extremely low but over time there is a tendency for attack to occur at weak points in the oxide film

Protecting aluminium by anodizing it

A surface treatment that artificially thickens the oxide film to make it even more protective

The film can even be coloured for decorative purposes, by adding colouring agents to the bath toward the end of the process

Standard Electrode Potentials

• The voltage when the metal is in equilibrium with a solution of its ions having a concentration of 1 mol per liter

• The same voltages as those shown in the Pourbaix diagrams for the

horizontal lines at the top of the immunity field when there is no

oxide film

Standard Electrode Potentials

The driving force for corrosion can be estimated by subtracting the potential for the metal from that for the oxygen reduction reaction.

Gold, platinum, and palladium are not corroded as there is no voltage driving force.

Corrosion Type

Pitting Corrosion

Preferential attack starts at

breaks or weaknesses in the

oxide film

Intergranular attack and stress corrosion cracking

Intergranular attack can occur when the grain boundaries have a lower corrosion resistance than the grains themselves.

stress corrosion cracking, a critical combination of

stress, material, and corrosive environment can lead to

cracks forming and growing

under static stress

Protecting ships’ hulls from corrosion

Magnesium sacrificial anode bolted to the steel hull just above the keel to prevent corrosion

Sacrificial protection (Cathodic protection)

The magnesium becomes the anode and corrodes. The iron becomes the cathode, where the oxygen reduction reaction takes place, and does not corrode.

Typical materials used are Mg, Al, and Zn.

Corrosion in reinforced concrete

Spalled reinforced concrete

Corrosion in reinforced concrete

Volume of rust produced is greater than the volume of steel that is lost.

• density of iron oxide is less than that of steel

• rust deposits are hydrated (FeO•H₂O) and full of

voids, their density is usually much less again

So if the steel reinforcement bars inside the concrete rust, they “expand,”

and crack open the surrounding concrete.

This occurs when there is not enough thickness of concrete over the

steelwork (“cover”) to protect it from the environment, or when there

is a lot of chloride ion around (immersion in seawater or salt-water

spray in the air)