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Electrochemical Corrosion

The conditions needed to promote many types of corrosion can be found in most oil wells. The basic electrochemical reactions, which occur simultaneously at the cathodic and anodic areas of metal and cause many forms of corrosion damage, are as follows:

1. At the cathode, the hydrogen (or acid) ion (H+) removes electrons from the cathodic. surfaces to form hydrogen gas (H2):

2e_+2H+ —> 2H° —> H2 (in acidic solution).

If oxygen is present, electrons are removed from the metal by reduction of oxygen:

4e"+02 + 4H+ —> 2 H20 (in acidic solution)

4 e~ + 2 H20 + 0— 2 —* 4 OH" (in neutral or alkaline solution)

2. At the anode, a metal ion (e. g., Fe2+) is released from its structural position in the metal through the loss of the bonding electrons and passes into solution in the water as soluble iron, or reacts with another component of the environment to form scale. The principal reaction is:

Fe — 2e" —> Fe2+

Thermodynamic data indicates that the corrosion process in many environments of interest should proceed at very high rates of reaction. Fortunately, experience shows that the corrosion process behaves differently. Studies have shown that as the process proceeds, an increase in concentration of the corrosion products develops rapidly at the cathodic and anodic areas. These products at the metal surfaces serve as barriers that tend to retard the corrosion rate. The reacting components of the environment may be depleted locally, which further tends to reduce the total corrosion rate.

The potential differences between the cathodic and anodic areas decrease as cor­rosion proceeds. This reduction in potential difference between the electrodes upon current flow is termed polarization. The potential of the anodic reaction approaches that of the cathode and the potential of the cathodic reaction ap­proaches that of the anode.

Electrode polarization by corrosion is caused by: changing the surface concentra­tion of metal ions, adsorption of hydrogen at cathodic areas, discharge of hydroxyl ions at anodes, or increasing the resistance of the electrolyte and films of metal — reaction products on the metal surface. Changes (increase or decrease) in the amount of these resistances by the introduction of materials or electrical energy into the system will change the corrosion currents and corrosion rate.

A practical method to control corrosion is through cathodic protection, whereby polarization of the structure to be protected is accomplished by supplying an external current to the corroding metal. Polarization of the cathode is forced beyond the corrosion potential. The effect of the external current is to eliminate the potential differences between anodic and cathodic areas on the corroding metal. Removal of the potential differences stops local corrosion action. Cathodic protection operates most efficiently in systems under cathodic control, i. e.. where cathodic reactions control the corrosion rate.

Materials may cause an increase in polarization and retard corrosion by absorbing on the surface of the metals and thereby changing the nature of the surface. Such materials act as inhibitors to the corrosion process. On the other hand, some materials may reduce polarization and assist corrosion. These materials, called depolarizers, either assist or replace the original reactions and prevent the buildup of original reaction products.

Oxygen is the principal depolarizer which aids corrosion in the destruction of metal. Oxygen tends to reduce the polarization or resistance, which normally develops at the cathodic areas, with the accumulation of hydrogen at these elec­trodes. The cathodic reaction with hydrogen ion is replaced by a reaction in which electrons at the cathodic areas are removed by oxygen and water to form hydroxyl ions (OH-) or water:

02 + 2H20 + 4e — —> 4 OH- (in neutral and alkaline solutions)

02 + 4H++4e — —» 2H20 (in acid solutions)

Polarization of an electrode surface reduces the total current and corrosion rate. Though the rate of metal loss is reduced by polarization, casing failures may increase if incomplete polarization occurs at the anodes. For example, inadequate anodic corrosion inhibitor will reduce the effective areas of the anodic surfaces and thus localize the loss of metal at the remaining anodes. This will result in severe pitting and the destruction of metal.

Resistances to the corrosion process generally do not develop to the same degree at the anodic and cathodic areas. These resistances reduce the corrosion rate, which is controlled by the slowest step in the corrosion process. Electrochemical corrosion comprises a series of reactions and material transport to and from the metal surfaces. Complete understanding of corrosion and corrosion control in a particular environment requires knowledge of each reaction which occurs at the anodic and cathodic areas.

Components of Electrochemical Corrosion

The various components which are involved in the process of corrosion of metal are: the metal, the films of hydrogen gas and metal corrosion products, liquid and gaseous environment, and the several interferes between these components.

Metal is a composite of atoms which are arranged in a symmetrical lattice struc­ture. These atoms may be considered as particles which are held in an ordered arrangement in a lattice structure by bonding electrons. These electrons, which are in constant movement about the charged particles, move readily throughout the lattice structure of metal when an electric potential is applied to the system. If bonding electrons are removed from their orbit about the particle center, the resulting cation will no longer be held in the metal’s crystalline structure and can enter the electrolyte solution.

Electrochemical corrosion is simply the process of freeing these cations from their organized lattice structure by the removal of the bonding electrons. Inasmuch as certain of the lattice electrons move readily within the metal under the influence of electrical potentials, the locations on the surface of the metal from which the cations escape and the locations from which the electrons are removed from the metal need not be and generally are not the same. Corrosion will not occur unless electrons are removed from some portion of the metal structure.

All metals are polycrystalline with each crystal having a random orientation with respect to the next crystal. The metal atoms in each crystal are oriented in a crystal lattice in a consistent pattern. 7’he pattern gives rise to differences in spacing and, therefore, differences in cohesive energy between the particles, which may cause preferred corrosion attack. Al the crystal boundaries the lattices are distorted, giving rise to preferred corrosion attack. In the manufacture and processing of metals, in order to gain desirable physical properties, both the composition and shape of the crystals may be made non-uniform, distorted or preferably oriented.

This may increase the susceptibility of the metal to corrosion attack. Tndistorted single crystals of metals experience comparatively little or no corrosion under the same conditions which may destroy commercial pieces of the same metal. Compositional changes in metal alloy crystals and crystal boundaries, which are present in steels and alloys, can promote highly localized corrosion.

Chemistry of Corrosion and Electromotive Force Series

Oxidation takes place when a given substance loses electrons or share of its elec­trons. On the other hand, reduction occurs when there is a gain in electrons by a substance. A substance that yields electrons to something else is called a reducing agent, whereas the substance which gains electrons is termed an oxidizing agent. Thus, electrons are always transferred from the reducing agent to the oxidizing agent. In the example below, two electrons are transferred from metallic iron to cupric ion:

Electrode reaction

Standard electrode potential. E° in Volts. 25°С

Li = Li+ + e"


K = K+ + e"


Ca = Ca2+ + 2e_


Na = Na+ + e~


Mg = Mg2+ + 2e~


Be = Be2+ + 2e_

+ 1.85

Al = Al3+ + 3e“

+ 1.67

Mn = Mn2+ + 2e“

+ 1.029

Zn = Zn2+ + 2e_


Cr = Cr3+ + 3e“


Ga = Ga3+ + 3e“


Fe = Fe2+ + 2e_


Cd = Cd2+ + 2e-


In = In3+ + 3e“


T1 = T1+ + e“


Co = Co2+ + 2e-


Ni = Ni2+ + 2e“


Sn = Sn2+ + 2e_


Pb = Pb2+ + 2e“


H2 = 2H+ + 2e-


Cu = Cu2+ + 2e“

— 0.345

Cu = Cu+ + e~

— 0.522

2Hg = Hg2+ + 2e-

— 0.789

Ag = Ag+ + e"

— 0.800

Pd = Pd2+ + 2e"

— 0.987

Hg = Hg2+ + 2e“

— 0.854

Pt = Pt2+ + 2e“

ca. — 1.2

Au = Au3+ + 3e“

— 1.50

Au = Au+ + e“

— 1.68

Fe° _ + Cu2+ -+ Fe2+ + Cu°

metallic cupric ferrous metallic

iron ion ion copper

The emf (electromotive force) series is presented in Table 6.1: potentials given are those between the elements in their standard state at 25°C and their ions at unit activity in the solution at 25°С. A plus sign ( + ) for E° shows that, for the above conditions, the reduced form of the reactant is a better reducing agent than H2. On the other hand, a negative (-) sign indicates that the oxidized form of the reactant is better oxidizing agent than H+. Thus, in general, any ion is better oxidizing agent than the ions above it.

Actual Electrode Potentials

In the emf series, each metal will reduce (or displace from solution) the ion of any metal below it in the series, providing all of the materials have unit activities. The activity of a pure metal in contact with a solution does not change with the environment. The activity of an ion, however, changes with concentration and the activity of a gas changes with partial pressure.

An electrode reaction, in which a metal M is oxidized to its ion Mn+, liberating n electrons, may be represented by the relation: M = Mn+ + nc~. The actual electrode potential of this reaction may be calculated from the standard elect rode potential by the use of the following expression:


E = E° — — In (АГ+) where:

E = actual electrode potential at the given concentration (Volts).

E° — standard electrode potential (Volts).

R — universal gas constant; 8.315 Volt Coulombs/°K.

T = absolute temperature (°K).

n = number of electrons transferred.

F — the Faraday, 96,500 Coulombs.

Mn+ = concentration of metal ions.

At 25°C (298°K), the formula becomes:

E = E°- log10 (M-)

77 4 /

The actual electrode potential for a given environment may be computed from the above relation. Table 6.2 shows how the actual electrode potentials of iron and cadmium vary with change in concentration of the ions.

Table 6.2: Variation in actual electrode potentials of iron and cadmium with change in concentration of ions.


Activity (moles/kg water)

1 0.1 0.01


Actual electrode potential


Fe = Fe2+ + 2e“

+0.440 +0.470 +0.499


C’d = C’d2+ + 2e“

+0.402 +0.431 +0.461


It is apparent from Table 6.2 that iron will reduce cadmium when their ion concentrations are equal, but the reverse holds true when the concentration of cadmium ion becomes sufficiently lower than that of the ferrous ion.

It is well to note that the standard electrode potentials are a part of the more general standard oxidation-reduction potentials. Textbooks on physical chemistry also contain a general expression for calculating the actual oxidation-reduction potential from the standard oxidation-reduction potential.

Galvanic Series

Dissimilar metals exposed to electrolytes exhibit different potentials or tendencies to go into solution or react with the environment. This behavior is recorded in tabulations in which metals and alloys are listed in order of increasing resistance to corrosion in a particular environment. Coupling of dissimilar metals in an electrolyte will cause destruction of the more reactive metal, which acts as an anode and provides protection for the less reactive metal, which acts as a cathode.

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