16-08-2012, 03:16 PM
Basics of Corrosion Chemistry
[attachment=31705]
Introduction
Metallic materials in practical use are normally exposed to corrosion in the
atmospheric and aqueous environments.Metallic corrosion is one of the problems
we have often encountered in our industrialized society; hence it has been studied
comprehensively since the industrial revolution in the late eighteenth century.
Modern corrosion science was set off in the early twentieth century with the
local cell model proposed by Evans [1] and the corrosion potential model proved by
Wagner and Traud [2]. The two models have joined into themodern electrochemical
theory of corrosion,which describesmetallic corrosion as a coupled electrochemical
reaction consisting of anodic metal oxidation and cathodic oxidant reduction. The
electrochemical theory is applicable not only to wet corrosion of metals at normal
temperature but also to dry oxidation of metals at high temperature [3].
Potential-pH Diagram
Thermodynamics shows that an electrode reaction is reversible at its equilibrium
potential, where no net reaction current is observed.We then learn that the anodic
reaction of metallic corrosion may occur only in the potential range more positive
than its equilibrium potential and that the cathodic reaction of oxidant reduction
may occur only in the potential range more negative than its equilibrium potential.
Moreover, it is known that metallic corrosion in aqueous solution is dependent not
only on the electrode potential but also on the acidity and basicity of the solution,
that is, the solution pH.
The thermodynamic prediction of metallic corrosion was thus illustrated by
Pourbaix [4] in the form of potential–pH diagrams, as shown for iron corrosion
in Figure 1.1. The corrosion of metallic iron may occur in the potential–pH
region where hydrated ferrous ions Fe2+, ferric ions Fe3+, and hydroxo-ferrous
ions Fe(OH)−
3 are stable. No iron corrosion occurs in the region where metallic
iron is thermodynamically stable at relatively negative electrode potentials. In the
regions where solid iron oxides and hydroxides are stable, no iron corrosion into
water is expected to develop and the iron surface is covered with solid oxide
films. In the diagram, we also see the equilibrium potentials of the hydrogen and
oxygen electrode reactions.
Corrosion Potential
An electrode of metal corroding in aqueous solution has an electrode potential,
which is called the corrosion potential. As a matter of course, the corrosion potential
stands somewhere in the range between the equilibrium potential of the anodic
metal dissolution and that of the cathodic oxidant reduction. It comes from the
kinetics of metallic corrosion that at the corrosion potential, the anodic oxidation
current of the metal dissolution is equal to the cathodic reduction current of the
oxidant. The corrosion kinetics is usually described by the electrode potential versus
reaction current curves of both the anodic oxidation and the cathodic reduction,
as schematically shown in Figure 1.2, which electrochemists call the polarization
curves of corrosion reactions. We see in Figure 1.2 that the intersecting point of the
anodic and cathodic polarization curves represents the state of corrosion, namely,
the corrosion potential and the corrosion current.
Passivity of Metals
Metallic passivity was discovered in 1790 by Keir [10], who found that metallic iron
violently corroding in the active state in concentrated nitric acid solution suddenly
turned into the passive state where almost no corrosion was observed. It was not
until 1960s that we confirmed the presence of an oxide film several nanometers
thick on the surface of passivatedmetals [11]. Latest overviews on metallic passivity
may be referred to in the literature of corrosion science [12].
We illustrate metallic passivity with the potential–current curve of anodic metal
dissolution for metallic iron, nickel, and chromium in acid solution as shown
in Figure 1.4. Anodic metal passivation occurs at a certain potential, called the
passivation potential, EP, beyond which the anodic current of metal dissolution
drastically decreases to a negligible level. It is an observed fact that the passivation
potential depends on the solution acidity, lineally shifting in the more positive
direction with decreasing solution pH. This fact thermodynamically suggests that
metallic passivity is caused by the formation of an oxide film on the metal, which
is extremely thin and invisible to the naked eye.
[attachment=31705]
Introduction
Metallic materials in practical use are normally exposed to corrosion in the
atmospheric and aqueous environments.Metallic corrosion is one of the problems
we have often encountered in our industrialized society; hence it has been studied
comprehensively since the industrial revolution in the late eighteenth century.
Modern corrosion science was set off in the early twentieth century with the
local cell model proposed by Evans [1] and the corrosion potential model proved by
Wagner and Traud [2]. The two models have joined into themodern electrochemical
theory of corrosion,which describesmetallic corrosion as a coupled electrochemical
reaction consisting of anodic metal oxidation and cathodic oxidant reduction. The
electrochemical theory is applicable not only to wet corrosion of metals at normal
temperature but also to dry oxidation of metals at high temperature [3].
Potential-pH Diagram
Thermodynamics shows that an electrode reaction is reversible at its equilibrium
potential, where no net reaction current is observed.We then learn that the anodic
reaction of metallic corrosion may occur only in the potential range more positive
than its equilibrium potential and that the cathodic reaction of oxidant reduction
may occur only in the potential range more negative than its equilibrium potential.
Moreover, it is known that metallic corrosion in aqueous solution is dependent not
only on the electrode potential but also on the acidity and basicity of the solution,
that is, the solution pH.
The thermodynamic prediction of metallic corrosion was thus illustrated by
Pourbaix [4] in the form of potential–pH diagrams, as shown for iron corrosion
in Figure 1.1. The corrosion of metallic iron may occur in the potential–pH
region where hydrated ferrous ions Fe2+, ferric ions Fe3+, and hydroxo-ferrous
ions Fe(OH)−
3 are stable. No iron corrosion occurs in the region where metallic
iron is thermodynamically stable at relatively negative electrode potentials. In the
regions where solid iron oxides and hydroxides are stable, no iron corrosion into
water is expected to develop and the iron surface is covered with solid oxide
films. In the diagram, we also see the equilibrium potentials of the hydrogen and
oxygen electrode reactions.
Corrosion Potential
An electrode of metal corroding in aqueous solution has an electrode potential,
which is called the corrosion potential. As a matter of course, the corrosion potential
stands somewhere in the range between the equilibrium potential of the anodic
metal dissolution and that of the cathodic oxidant reduction. It comes from the
kinetics of metallic corrosion that at the corrosion potential, the anodic oxidation
current of the metal dissolution is equal to the cathodic reduction current of the
oxidant. The corrosion kinetics is usually described by the electrode potential versus
reaction current curves of both the anodic oxidation and the cathodic reduction,
as schematically shown in Figure 1.2, which electrochemists call the polarization
curves of corrosion reactions. We see in Figure 1.2 that the intersecting point of the
anodic and cathodic polarization curves represents the state of corrosion, namely,
the corrosion potential and the corrosion current.
Passivity of Metals
Metallic passivity was discovered in 1790 by Keir [10], who found that metallic iron
violently corroding in the active state in concentrated nitric acid solution suddenly
turned into the passive state where almost no corrosion was observed. It was not
until 1960s that we confirmed the presence of an oxide film several nanometers
thick on the surface of passivatedmetals [11]. Latest overviews on metallic passivity
may be referred to in the literature of corrosion science [12].
We illustrate metallic passivity with the potential–current curve of anodic metal
dissolution for metallic iron, nickel, and chromium in acid solution as shown
in Figure 1.4. Anodic metal passivation occurs at a certain potential, called the
passivation potential, EP, beyond which the anodic current of metal dissolution
drastically decreases to a negligible level. It is an observed fact that the passivation
potential depends on the solution acidity, lineally shifting in the more positive
direction with decreasing solution pH. This fact thermodynamically suggests that
metallic passivity is caused by the formation of an oxide film on the metal, which
is extremely thin and invisible to the naked eye.