THERMODYNAMICS OF CORROSION
For an anodic oxidation reaction of metal to occur, a simultaneous
reduction must take place. In corroding metal systems, the anodic and
cathodic half-reactions are mutually dependent and form a galvanic or
spontaneous cell. A cell in which electrons are driven by an external
energy source in the direction counter to the spontaneous half-reactions
is termed an electrolytic cell. A metal establishes a potential or emf with
respect to its environment and is dependent on the ionic strength and
composition of the electrolyte, the temperature, the metal or alloy itself,
and other factors. The potential of a metal at the anode in solution arises
from the release of positively charged metal cations together with the
creation of a negatively charged metal. The standard potential of a metal
is defined by fixing the equilibrium concentration of its ions at unit
activity and under reversible (zero net current) and standard conditions
(1 atm, 101 kPa; 25°C). At equilibrium, the net current density
(mA/cm2) of an electrochemical reaction is zero. The nonzero anodic
and cathodic currents are equal and opposite, and the absolute magnitude
of either current at equilibrium is equal to the exchange current
density (that is, ian 5 icath 5 io ).
The potential, a measure of the driving influence of an electrochemical
reaction, cannot be evaluated in absolute terms, but is determined by
the difference between it and another reference electrode. Common
standard reference electrodes used are saturated calomel electrode (SCE,
0.2416 V) and saturated Cu/CuSO4 (0.337 V) whose potentials are
measured relative to standard hydrogen electrode (SHE), which by definition
is 0.000 V under standard conditions. Positive electrochemical
potentials of half-cell reactions versus the SHE (written as a reduction
reaction) are more easily reduced and noble, while negative values signify
half-cell reactions that are more difficult to reduce and, conversely,
more easily oxidized or active than the SHE. The signs of the potentials
are reversed if the SHE and half-cell reactions of interest are written as
oxidation reactions. The potential of a galvanic cell is the sum of the
potentials of the anodic and cathodic half-cells in the environment
surrounding them.
From thermodynamics, the potential of an electrochemical reaction is
a measure of the Gibbs free energy DG52nFE, where n is the number
of electrons participating in the reaction, F is Faraday’s constant
(96,480 C/mol), and E is the electrode potential. The potential of the
galvanic cell will depend on the concentrations of the reactants and
products of the respective partial reactions, and on the pH in aqueous
solutions. The potential can be related to the Gibbs free energy by the
Nernst equation
DE 5 DE° 1
2.3RT
nF
log
(ox)x
(red)r
where DE° is the standard electrode potential, (ox) is the activity of an
oxidized species, (red) is the activity of the reduced species, and x and r
are stoichiometric coefficients involved in the respective half-cell reactions.
Corrosion will not occur unless the spontaneous direction of the
reaction (that is, DG , 0) indicates metal oxidation. The application of
thermodynamics to corrosion phenomena has been generalized by use
of potential-pH plots or Pourbaix diagrams. Such diagrams are constructed
from calculations based on the Nernst equation and solubility
data for various metal compounds. As shown in Fig. 6.5.1, it is possible
to differentiate regions of potential versus pH in which iron either is
immune or will potentially passivate from regions in which corrosion
will thermodynamically occur. The main uses of these diagrams are to:
(1) predict the spontaneous directions of reactions, (2) estimate the
composition of corrosion products, and (3) predict environmental
changes that will prevent or reduce corrosion. The major limitations of
Pourbaix diagrams are that (1) only pure metals, and not alloys, are
usually considered; (2) pH is assumed to remain constant, whereas HER
may alter the pH; (3) they do not provide information on metastable films; (4) possible aggressive solutions containing Cl2, Br2, I2, or
NH4 1 are not usually considered; and (5) they do not predict the kinetics
or corrosion rates of the different electrochemical reactions. The predictions
possible are based on the metal, solution, and temperature designated
in the diagrams. Higher-temperature potential-pH diagrams have been developed (Computer-Calcuated Potential-pH Diagrams to
300°C, NP-3137, vol. 2. Electric Power Research Institute, Palo Alto,
CA). Ellingham diagrams provide thermodynamically derived data for
pure metals in gaseous environments to predict stable phases, although
they also do not predict the kinetics of these reactions.

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