LECTURE № 13 CORROSION 15.04.2015

chemical reaction with their environment.
In the most common use of

the word, this means electrochemical oxidation of metals in reaction with an oxidants such as oxygen. Rusting, the formation of iron oxides, is a well-known example of electrochemical corrosion. This type of damage typically produces oxide(s) or salt(s) of the original metal.

as ceramics or polymers, although in this context, the term

degradation is more common. Corrosion degrades the useful properties of materials and structures including strength, appearance and permeability to liquids and gases.

air, but the process can be strongly affected by exposure

to certain substances. Corrosion can be concentrated locally to form a pit or crack, or it can extend across a wide area more or less uniformly corroding the surface. Because corrosion is a diffusion-controlled process, it occurs on exposed surfaces. As a result, methods to reduce the activity of the exposed surface, such as passivation and chromate conversion, can increase a material's corrosion resistance. However, some corrosion mechanisms are less visible and less predictable.

as a result of exposure to chemical and electrochemical environment,

it is heterogeneous redox process that occurs at the interface.

reduced to the oxidation of the metal and restore oxidant.

When metal corrosion occur at its surface take place simultaneously two independent electrochemical reactions:
at anode: Me - ne = Me+n
at cathode: Ox + ne = Red

Corrosive environment called surfactants, are present around the structural member, its impact on the material and cause it to corrode. Corrosive medium may be air, industrial atmosphere, gases, water, sea climate, land - soil, acids, alkalis, water and salt solutions.

of iron in aqueous solution containing hydrogen ions (acid medium,

pH<7):
2Fe + 6H+ = 2Fe3+ + 3H2

Corrosion of iron in water containing oxygen (neutral medium, pH=7):
4Fe + 3O2 + 2H2О = 2Fe2O3·H2О

the types of corrosive process
Chemical corrosion
Electrochemical corrosion

continual


local


intergranual

gas
liquid
soil
atmospheric
stray currents

corrosion
contact corrosion
corrosion fatigue
corrosion
cracking
Corrosion friction
(Fretting)
Corrosion in the collision
(Cavitation)

the two metals, which enters the electrolyte.
Contact corrosion - electrochemical

corrosion of any two metals with different potentials in contact, dipped in an electrolyte.

high temperatures (e.g., in the combustion chamber jet).
Atmospheric corrosion -

corrosion in air (in the presence of a condensed film of moisture and precipitation).
Liquid corrosion - the chemical destruction of metal in fluids, electrolytes which are not.

Soil corrosion - corrosion of metal products in contact with the soil - soil electrolyte.
Stray currents corrosion is destruction of metal structures due to the ingress of corrosion of the conductive medium (soil, sea water) electric (so called stray) current. Source stray currents in the soil often electrified railways, as well as current generators (eg, welders), grounded on the ground.

in which the oxidation of the metal and restoring the

oxidant corrosive environment occurs in one act without causing an electric current.
This type of corrosion occurs due to direct chemical attack of environment or atmospheric gases like oxygen halogens, hydrogen sulphide, sulphur dioxide, nitrogen or anhydrous inorganic liquid with metal surfaces in immediate proximity.
Chemical corrosion is of three main types as:
Oxidation Corrosion (with Oxygen)
Corrosion by Other Gases
In non-electrolyte Liquid Metal Corrosion

with oxygen in the absence of moisture:


Alkali and alkaline

earth metals react with oxygen at room temperature and form corresponding oxides, while some metals react with oxygen at higher temperature.

Metals like Ag, Au and Pt are not oxidized as they are noble metals.

4Fe + 3О2 = 2Fe2О3

as a thin film on the metallic surface which protects

the metal from further corrosion:



If diffusion of either oxygen or metal is across this layer, further corrosion is possible. Thus, the layer of metal oxide plays an important role in the process of corrosion.

inhibit further corrosion. They form a stable, tightly adhering oxide

film.
In case of porous oxide film, atmospheric gases pass through the pores and react with the metal and the process of corrosion continues to occur till the entire metal is converted into oxide.
Porous oxide layer is formed by alkali and alkaline earth metals. Molybdenum forms a volatile oxide film of MoO3 which accelerates corrosion.
Au, Ag, Pt form unstable oxide layer which decomposes soon after the formation, thereby preventing further corrosion.

atmosphere, these gases react with metal and form corrosion products

which may be protective or non-protective.

Dry Cl2 reacts with Ag and forms AgCl which is a protective layer, while SnCl4 is volatile.

In petroleum industries at high temperatures, H2S attacks steel forming FeS scale which is porous and interferes with normal operations.

2Fe + 3Cl2 = 2FeCl3

through metallic pipes and causes corrosion due to dissolution or

due to internal penetration.
For example:
liquid metal mercury dissolves most metals by forming amalgams, thereby corroding them.
In oil
 In sulphur
 In organic substances

Cu + S = CuS

2 Ag + S = Ag2S

2Al + 6ССl4 = 3C2Cl6 + 3AlCl3

elements in which the oxidation of the metal and restoring

the oxidant corrosive environment is a result of the flow of several elementary acts, accompanied by the appearance of a galvanic couple and electric current.
This type of corrosion occurs when the metal comes in contact with a conducting liquid or when two dissimilar metals are immersed or dipped partly in a solution. There is the formation of a galvanic cell on the surface of metals. Some parts of the metal surface act as anode and rest act as cathode. Water must be present to serve as a medium for the transport of ions.
The most common depolarizers are oxygen, acids, and the cations of less active metals.

corrosion at anode, while reduction takes place at cathode:









The corrosion

product is formed on the surface of the metal between anode and cathode.

At cathode:

Hydrogen depolarization

Oxygen depolarization

and catholic regions. The anodic reaction involves dissolution of the

metal as corresponding metallic tops with the liberation of tree electrons.
Anodic area:  Me-ne- = Me+n                       (Oxidation)
Cathodic area: Reduction reaction consumes electrons and depending on the nature, of corrosive environment the following reactions may take place:
Hydrogen evolution: 2H+ + 2e- = H2
Oxygen reduction (acidic medium pH<7):
           O2 + 4H+ + 4e- = 2H2O
Oxygen reduction (neutral or alkaline medium):
           O2 + 2H2O + 4e- = 4OH-
Metal ion reduction:   M3+ + e- = M2+
Metal deposition:       M+ + e- = M

solutions) is in contact with a metal or
when two dissimilar,

metals or alloys are either immersed or dipped partially in a aqueous solution.

In electrochemical or wet corrosion the following observations are made:
Formation of anodic and catholic part or parts in contact with each, other.
Presence of conducting medium
Corrosion of anodic areas only.
Formation of corrosion product between anodic and catholic areas.

* nH2O
Oxygen reduction (neutral or alkaline medium):
           2Fe + O2 +

3H2O = 2Fe(OH)3

– 2e = Fe2+

At cathode: 2H+ + 2e

= H2

Fe + H2SO4 = FeSO4 + H2

galvanized steel

At anode: Zn – 2e = Zn2+

At cathode: Fe+2 + 2e = Fe

Zn + H2SO4 = ZnSO4 + H2

the metal and the environment. The important factors which may

influence the corrosion process are:
Nature of the metal (position Me in ECS)
Environment (oxidant type)
Concentration of electrolyte
Temperature
Hydrogen over voltage (pH)

0.00
Copper
Silver
Platinum
Gold---------------------- +1.15 V
Metal

SRP, Eo

Decreasing tendency
to loose
electrons

Increasing order of std reduction potential

Less Reduction

More oxidation

of corrosion with oxygen depolarization possible, example Cu (0,34 В)
area

of corrosion with hydrogen depolarization possible, example Fe (-0,44 B)

-0,41 B

+ 0,82 В

rise in temperature of environment.
Humidity of the air: Rate

of corrosion increases with presence of moisture in the atmosphere.
Impurities: Presence of impurities like CO2, H2S, SO2, acid fumes etc increases corrosion rate.
Influence of pH: In acid medium corrosion is more and in alkaline medium it is less.

group of the sub-groups - very resistant metals
Group 2 –

are unstable,
Group 2 sub-groups – are more stable (in the presence of oxygen to form a strong oxide film protecting from further destruction)

corrosion properties of metals

it is destroyed in solutions of acids and bases). In

concentrated nitric and sulfuric acids passivated aluminum.

Group 4 – tin (Sn) and lead (Pb) – corrosion-resistant metals, thanks to strong oxide film.

to passivate and hence high resistance to corrosion.

Osmium, iridium, platinum

- the most resistant to corrosion

Iron passivated by concentrated sulfuric and nitric acids

Cathodic or Anodic protection
Active electrochemical protection
Coatings

have been presented above. It is especially imortant that you

know the precise meanings of all the highlighted terms in the context of this topic.
Electrochemical corrosion of metals occurs when electrons from atoms at the surface of the metal are transferred to a suitable electron acceptor or depolarizer. Water must be present to serve as a medium for the transport of ions.
The most common depolarizers are oxygen, acids, and the cations of less active metals.
Because the electrons flow through the metallic object itself, the anodic and cathodic regions (the two halves of the electrochemical cell) can be at widely separated locations.
Anodic regions tend to develop at locations where the metal is stressed or is protected from oxygen.
Contact with a different kind of metal, either direct or indirect, can lead to corrosion of the more active one.
Corrosion of steel can be inhibited by galvanizing, that is, by coating it with zinc, a more active metal whose dissolution leaves a negative charge on the metal which inhibits the further dissolution of Fe2+.
Cathodic protection using an external voltage source is widely used to protect underground structures such as tanks, pipelines and piers. The source can be a sacrificial anode of zinc or aluminum, or a line-operated or photovoltaic power supply.