Corrosion Mechanisms
2.3 Potential–pH Diagrams
Potential–pH diagrams, also known as Pourbaix diagrams, are graphical rep-resentations of the stability of a metal and its corrosion products as a func-tion of the potential and pH (acidity or alkalinity) of the aqueous solufunc-tion.
The potential is shown on the vertical axis and the pH on the horizontal axis.
Such diagrams are constructed from calculations based on the Nernst equa-tion and the solubility data for various metal compounds. The potential–pH diagram for an Fe–Η2Ο system is shown in Figure 2.3. In the diagram, the horizontal lines represent pure electron transfer reactions dependent solely on potential, but independent of pH:
1. Fe Fe= 2++2e− (2.3)
2. Fe2+=Fe3++e− (2.4)
These lines extend across the diagram until the pH is sufficiently high to facilitate the formation of hydroxides, represented by vertical lines, thereby reducing the concentration Fe2+ and Fe3+ ions. The boundary is often set arbitrarily at the concentration of these ions at 10−6 g-ions/liter, which is indicative of a negligible dissolution or corrosion of the metal in the medium.
The vertical lines in Figure 2.3 correspond to the reactions:
3. Fe2++2H O Fe OH2 =
( )
2+2H+ (2.5)4. Fe3++3H O Fe OH2 =
( )
3+3H+ (2.6) There is no electron transfer involved and the reactions are solely dependent on pH.The sloping lines in Figure 2.3 represent equilibria involving both electron transfer and pH; for example:
Potential–pH (Pourbaix) diagram for Fe–H2O system.
Corrosion Mechanisms 13
The hydrogen and oxygen are also shown in the diagram by the dotted lines.
The hydrogen line represents the equilibria:
2H + 2e = H in acid solutions+ – 2 (2.9)
or
2H O2 +2e−=H2+2OH in neutral or alkaline solu− ttions (2.10)
These two reactions are equivalent and their pH dependence of single elec-trode potential is represented by:
EH /H+ EOH /H pH
2 +
= 2−0 059. (2.11)
at pH = 0; that is, for [H+] = 1, EOH+/H2=0 and the slope is −0.059V. Similarly, for oxygen equilibrium with water the corresponding reactions at lower and higher pH are:
O2+4H++4e−=2H O2 (2.12)
and
O2+2H O2 +4e−=4OH− (2.13)
The pH dependence of single electrode potential is represented by:
EO2/H O2 =EOO /H O2 2 −0 059. pH (2.14)
at pH = 0, EOO /H O2 2 =1 226. V and at pH = 1 (i.e., for [OH−] = 1), EOO /H O2 2 =0 401. V. Here again, the slope of the line is −0.059V. Water is stable in the area designated by these two lines. Below the hydrogen line it is reduced to hydrogen gas, and above the oxygen line it is oxidized to oxygen.
The potential–pH diagram shows three clear-cut zones:
1. Immunity zone. Under these conditions of potential and pH, iron remains in metallic form.
2. Corrosion zone. Under these conditions of potential and pH, iron cor-rodes, forming Fe2+ or Fe3+ or HFeO2−.
3. Passive zone. Under these conditions of potential and pH, protective layers of Fe(OH)3 form on iron and further corrosion of iron does not take place.
Such diagrams can be used for:
1. Predicting the spontaneous direction of reactions
2. Estimating the stability and composition of corrosion products 3. Predicting environmental changes that will prevent or reduce
corrosion
With reference to Figure 2.3, corrosion prevention can be achieved by lower-ing the electrode potential down to the zone of immunity, raislower-ing the elec-trode potential up to the region of passivity, or raising the pH or alkalinity of the solution so that a passive film is formed.
There are, however, limitations in using such diagrams. The most impor-tant of these is that they represent equilibrium conditions and hence cannot be used for predicting the rate of a reaction. The tacit assumption that cor-rosion products (oxides, hydroxides, etc.) lead to passivity may not always be true because they may not always precipitate on the metal surface. The possibility of precipitation of other ions such as chlorides, sulfates, and phos-phates has been ignored. Finally, the pH at the metal surface may vary dras-tically because of side reactions, and a prediction of corrosion based on the bulk pH of the solution may be misleading.
2.4 Polarization
At an intermediate resistance in the circuit, some current begins to flow and the potentials of both half-cell reactions move slightly toward each other.
This change in potential is called polarization. The resistance in the circuit depends on a number of factors, including the resistivity of the media, surface films, and the metal itself. The relationships between polarization reactions at each half-cell are represented in Figure 2.4. The intersection of the two polarization lines (curves) closely approximates the corrosion cur-rent and the combined cell potentials for the freely corroding situation.
Once the corrosion current is determined, the corrosion density can be calculated by determining the surface area. Using Faraday’s laws, a corro-sion rate in terms of metal loss per unit time can be determined. However, polarization data can be more useful than just estimating corrosion rates.
The extent of polarization can help predict the type and severity of cor-rosion. As polarization increases, corrosion decreases. Polarization may
Corrosion Mechanisms 15
be preferable to either cathodic or anodic reactions. Understanding the influence of environmental changes on polarization can offer insight into controlling corrosion. For example, in the iron–hydrochloric acid exam-ple, hydrogen gas formation at the cathode can actually slow the reaction (increased current resistance) by blocking access of hydrogen ions to the cathode site. This results in cathodic polarization and lowers the current flow and corrosion rate. If oxygen is bubbled through the solution, the hydrogen is removed more rapidly by combining to form water and the corrosion rate increases significantly. Although this is an oversimplified view of the effects of oxygen, it does indicate that the degree of polar-ization can be affected by changes in the environment, either natural or induced.
There are three basic causes of polarization, termed activation, concentra-tion, and potential drop. Potential drop is the change in voltage associated with effects of the environment and the circuit between the anode and cath-ode sites. It includes the effects of the resistivity of the media, surface films, corrosion products, etc.
2.4.1 activation Polarization
Activation polarization arises out of a slow step in the electrode reaction for which an activation energy in the form of an increment in potential is required for the reaction to proceed. This will best be illustrated by the hydrogen evolution reaction.
The hydrogen evolution reaction consists of several steps as shown in Figure 2.5. Either the electron transfer step (step 2) or the formation of
Current, i icorr
2H+ + 2e H2
Fe Fe2+ + 2e
–0.5 –0.3 –0.1
Potential, E
E (H2/H+) Cathodic polarization curve 0.1
Ecorr
E (Fe/Fe2+) Anodic polarization curve
FigurE 2.4
Polarization of iron in acid.
hydrogen molecules (step 3) is deemed the slowest step in the reaction sequence, and the rate of overall reaction will depend on how fast or slow it proceeds. Therefore, to have a higher rate of reaction, expressed in terms of increased current density, an increase in potential should be effected. The relationship between reaction rate and change in potential (overvoltage) is expressed by the Tafel equation:
ηa β
o
log i
= ± i
(2.15) where ηa is overvoltage polarization (in volts), and β is a constant, called the Tafel constant (also expressed in volts), and is usually on the order of 0.1V.
A graphical representation of Equation 2.15, as applied to the hydrogen evolution reaction, with a β slope of 0.1V is shown in Figure 2.6. It can be noted from the graph that 0.1V change in overvoltage can effect a tenfold increase or decrease in the reaction rate.
Dissolution reactions (anodic) in corrosion are usually controlled by activation polarization where the solution of ions is the probable rate-trolling step. Hydrogen evolution reactions (cathodic reactions) are con-trolled by activation polarization when the concentration of hydrogen ions is high.
H+– H+
H –2
H
H+ H+
–H2 H2
H2
H2 H2 1
3
3
Zinc 2–
4
1
FigurE 2.5
Steps involved in hydrogen reduction reaction.
Corrosion Mechanisms 17
2.4.2 Concentration Polarization
A buildup or depletion of ions at the electrode surface as a result of reac-tion will change the value of the electrode potential according to the Nernst equation. For example, for a corroding zinc electrode, the concentration of zinc will increase with dissolution in the vicinity of the electrode. The value of aoxid in the equation increases, causing the electrode potential to shift in a positive direction.
For the hydrogen evolution reaction, the higher rate of discharge of hydro-gen ions at the electrode surface brings down the value of aoxid and the electrode potential, according to the Nernst equation, shifts in a negative direction. However, the rate of discharge of hydrogen ions at the electrode surface depends on the diffusion of hydrogen ions from the bulk of the solu-tion to the surface:
i = DnFC
l x (2.16)
100 10
1 0.1
–0.3 0.01 –0.2 –0.1 0 0.1
2H+
H2
H2 2H+ + 2e
–
2e– + NobleActiveOvervoltage ηa Volts
0.2
Current Density ioH2/H+
FigurE 2.6
Activation polarization curve of a hydrogen electrode.
where il is called the limiting diffusion current density (amp/cm2), D is the diffusion coefficient for Η+ ion, n is the number of electrons transferred, F is the Faraday number, C is the bulk concentration of Η+ ions in the solution, and x is the thickness of the diffusion layer adjacent to the electrode surface through which the concentration of the reacting species (H+ ions) changes from C in the bulk to zero at the electrode surface.
A mathematical expression for concentration polarization involves il and is given by:
where ηc is overvoltage due to concentration polarization (in volts). A graphi-cal representation of the equation is shown in Figure 2.7.
It can be seen from the graph in Figure 2.7 as well as from Equation 2.17 that as i approaches il, ηc tends to infinity. As evident from Equation 2.16, factors such as increasing velocity (smaller x), increasing temperature (higher D), and increasing concentrations will increase the value of iL, that is, a shift in the vertical position of the curve in Figure 2.7 more toward the right.
There is no question of concentration polarization when the supply of reacting species is abundant. Hence, in metal dissolution reactions, its effect is negligible as the supply of metal atoms for dissolution is unlimited. On the other hand, for a hydrogen evolution reaction, concentration polarization becomes significant in the solutions of low Η+ concentration. More often, the reduction process is controlled by a combined polarization — that is, activa-tion polarizaactiva-tion at lower reacactiva-tion rates and concentraactiva-tion polarizaactiva-tion at higher reaction rates — as i approaches il. A graphical representation of such combined polarization is shown in Figure 2.8.
Increasing velocity
(a) Concentration polarization curve for reduction process, and (b) effect of environmental variations on concentration polarization curve.
Corrosion Mechanisms 19