Corrosion Mechanisms
7
2.1 Cell Potentials
Understanding electrochemical behavior and possible reactions can help in
predicting the possibility and extent of corrosion. A reaction will occur only
if there is a negative free energy change (ΔG). For electrochemical reactions,
the free energy change is calculated from:
∆ =−
G
nFE
where n is the number of electrons, F is Faraday’s constant, and Ε is the
cell potential.
Therefore, for a given reaction to take place, the cell potential is
taken as the difference between the two half-cell reactions — the one
at the cathode minus the one at the anode. The half-cell potential exists
because of the difference in the neutral state compared to the oxidized
state, such as Fe/Fe 2+ , or at the cathode, the difference between the neutral state and the reduced state as in Η + /Η 2 . These reduction–oxidation
(redox) potentials are measured relative to a standard half-cell potential. The chart in Table 2.1 lists potentials relative to the Η + /Η 2, which is
set at zero.
Looking at the example of iron corroding freely in acid, the cell potential
is calculated as:
Ε = Cathode half-cell – Anode half-cell
Ε = E(H + /H 2 ) – E(Fe/Fe 2+ )
Ε = 0 − (0.440) = +0.44
Because the cell potential is positive, the reaction can take place.
The larger the potential difference, the greater the driving force for the
reaction. Whether corrosion does occur and at what rate depends on other
factors. For corrosion to occur, there must be a current flow and a completed
circuit, which is then governed by Ohm’s law: I = E/R. The cell potential
calculated here represents the peak value for the case of two independent
reactions. If the resistance were infinite, the cell potential would remain as
calculated but there would be no corrosion at all. If the resistance of the circuit is zero, the potentials of each half-cell would approach the other while
the rate of corrosion would be infinite.
7
2.1 Cell Potentials
Understanding electrochemical behavior and possible reactions can help in
predicting the possibility and extent of corrosion. A reaction will occur only
if there is a negative free energy change (ΔG). For electrochemical reactions,
the free energy change is calculated from:
∆ =−
G
nFE
where n is the number of electrons, F is Faraday’s constant, and Ε is the
cell potential.
Therefore, for a given reaction to take place, the cell potential is
taken as the difference between the two half-cell reactions — the one
at the cathode minus the one at the anode. The half-cell potential exists
because of the difference in the neutral state compared to the oxidized
state, such as Fe/Fe 2+ , or at the cathode, the difference between the neutral state and the reduced state as in Η + /Η 2 . These reduction–oxidation
(redox) potentials are measured relative to a standard half-cell potential. The chart in Table 2.1 lists potentials relative to the Η + /Η 2, which is
set at zero.
Looking at the example of iron corroding freely in acid, the cell potential
is calculated as:
Ε = Cathode half-cell – Anode half-cell
Ε = E(H + /H 2 ) – E(Fe/Fe 2+ )
Ε = 0 − (0.440) = +0.44
Because the cell potential is positive, the reaction can take place.
The larger the potential difference, the greater the driving force for the
reaction. Whether corrosion does occur and at what rate depends on other
factors. For corrosion to occur, there must be a current flow and a completed
circuit, which is then governed by Ohm’s law: I = E/R. The cell potential
calculated here represents the peak value for the case of two independent
reactions. If the resistance were infinite, the cell potential would remain as
calculated but there would be no corrosion at all. If the resistance of the circuit is zero, the potentials of each half-cell would approach the other while
the rate of corrosion would be infinite.