In this article we will discuss about:- 1. Introduction to Corrosion 2. Classification of Corrosion 3. Rate 4. Corrosion Rates from Polarisation and Mixed Potential Theory.
Introduction to Corrosion:
Corrosion is responsible for colossal loss of materials occurring everywhere every moment involving billions and billions of rupees annually. It is one of the most serious destructive agents, and is one of the major technological problems of modern society. Though a lot of scientific understanding of many phases of corrosion has been developed, but it is a very complex phenomenon as there are a large number of complex variables responsible for it.
Corrosion is the deterioration and loss of metal, or material by chemical, or electrochemical reactions with its environment. The definition includes all types of atmospheres, and at all temperatures. The commonly seen rusting of iron and steel is the best example of corrosion; whereas the regular painting of bridges, ships, house-hold structures, etc., is a step towards prevention or reduction in the rate of corrosion of them.
The steel structure of a pier near the sea is severely corroded because of its intermittent immersion in sea water by the usual tides, and the corrosion is much more than if it remains covered permanently in sea water; which is further aggravated by the presence of bacteria and sea-weeds.
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There are quite a few practical advantages of corrosion and oxidation- The corrosion product as a patina of bluish-green colour on ancient plaques and statues of copper, has the aesthetic value as it pleases the eye; the working of lead-acid battery depends on corrosion reactions; the chemical etching of metals and alloys is by corrosion; anodic oxidation is the basis of electrolytic polishing: electro-chemical machining and oxygen-flame cutting are useful for shaping metals; etc., but these are of hardly any significance as compared to innumerable and massive loss due to corrosion.
Corrosion is a very natural process, because in nature, most metallic elements (except noble metals) occur in the most stable state having low energy, normally as oxides (as metallic ores). Energy is used to convert these oxides to metallic state to be useful to mankind against this natural tendency.
Thus, the metals, or materials (except some noble metals) are in higher energy states, and try to revert to an oxidised state. The main reason of oxidation and corrosion, i.e. of reactions between metals and their environments is the tendency to decrease the energy of the system as a result of the corrosion reactions. The practical solution of the problems of corrosion, thus, lies in controlling it by reducing the rate at which the corrosion reactions proceed.
Corrosion has gained so much-importance that an engineer, now while selecting an alloy for an application, also takes into account the corrosion resistance of the alloy in the specific environment.
Classification of Corrosion:
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Though practically all corrosion is electro-chemical in nature but a better understanding of corrosion is developed by discussing it in three broad categories such as direct-chemical corrosion and electro-chemical corrosion.
In direct-chemical corrosion, the metal dissolves in a corrosive liquid medium. The metal continues to dissolve in the liquid until either the metal is consumed, or the liquid is saturated. For example, pickling of steel is commonly done (before plating, etc.) by immersing the steel in an acid pickling solution (5-10% H2SO4 in water at 60-65°C).
The acid attacks the metal and dissolves the surface uniformly to make the surface smooth with an etched appearance. No selective attack of a certain phase or component takes place and no protective layer is formed.
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Metal goes into solution as ferrous ions, and the hydrogen ions present in the solution absorb the electrons from the metal to form hydrogen gas bubbles on the steel surface as:
Fe + 2H+ → Fe++ + H2 (gas) …(14.1)
The electrical effect, in which hydrogen ions pick up the electrons from iron, takes place in the immediate vicinity of the chemical reaction. This differs from electro-chemical corrosion, where electrical current flows through parts of metal in which corrosion is not taking place, and the iron is dissolved at the anode and the hydrogen gas is evolved at the cathode.
Alkaline solutions also have direct-chemical corrosion on some metals like aluminium, tin, zinc, etc. The sensitivity of kitchen utensils made of aluminium to washing soda is familiar.
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Caustic soda reacts with metal such as zinc as:
Zn + 2 NaOH – Na2 ZnO2 + H2 (gas)
Chemical corrosion takes place at an almost constant but high rate, and can be measured as milligram per square decimeter per day (mdd).
Refractories, used as lining in metal melting or refining furnaces, may react with the slag produced during the process, requiring the replacement of lining quite frequently. This is direct-chemical attack. In the initial blow period of basic Bessemer converter for steel making, the silica produced may react with the basic lining of the furnace.
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In some cases, selective leaching may take place, i.e., one specific element may be selectively dissolved, or leached, from the solid. In gray-iron buried-gas-lines, graphitic corrosion may occur, i.e., iron is selectively dissolved in water, or soil leaving behind interconnected graphite flakes to cause leakage or even explosions.
Dezincification of brass having more than 15% Zn is another example. Though, here, copper and zinc are dissolved in aqueous solutions at high temperatures, but copper ions are replated on the brass and zinc ions stay back in the solution. Brass is then porous and weak.
It has been seen that small sized ions, or molecules dissolve faster than complex ions, and the reaction is faster at high temperatures.
It is a bit difficult to prevent direct-chemical corrosion. The best way to prevent is to avoid contact between the metal and that liquid such as by having coaling in between them. Tin plating on the steel surface prevents direct-corrosion by dilute acids.
2. Electro-Chemical Corrosion:
Electro-chemical corrosion, probably, is the most common mechanism by which corrosion occurs at or near room temperature, as a result of reaction of metals with water, or aqueous solutions of salts, acids, or bases. It is named electro-chemical corrosion as a chemical reaction accompanied with the passage of electric-current takes place. Although there are many variations of this type of corrosion, Fig. 14.10 illustrates its essential characteristics. This Fig. illustrates galvanic corrosion, which is one of the most common reasons of electro-chemical corrosion.
When two dissimilar metals such as iron and copper, are placed in electrical contact in the presence of an electrolyte, such as a salt solution having dissolved oxygen, a potential difference between iron and the copper is established. The potential of each metal is related to its relative tendency to dissolve.
One of the metals, here iron, having lower potential acts as anode, and dissolves (i.e., corrodes) by the reaction, which is always an oxidation reaction:
Anodic reaction, Fe → Fe2+ + 2 e̅ …(14.12)
The electrons, thus, liberated, migrate from the anode (iron) to the higher potential area, the cathode (here copper), where these are consumed in the cathodic reaction. In situation of Fig. 14.10 (a), hydroxyl ions, (OH) ̅ are produced by the following cathodic reaction. Other cathodic reactions also consume electrons.
Cathodic reactions are reduction reactions, and usually donot effect the cathode metal, since most metals cannot be reduced further:
Cathodic reaction, H2O + 2e̅ + 1/2 O2 (dissolved oxygen) → 2 (OH) ̅…(14.13)
The positively charged cations, Fe2+ migrate through the electrolyte towards the cathode, and the anions, (OH) ̅migrate towards the anode to produce an insoluble compound at a distance from the anode by the reaction:
Fe2+ + 2(OH) ̅ → Fe(OH)2 …(14.14)
A current flows through the circuit formed by the metals and the solution; the electrons are the current carriers in the metals, whereas ions in the solution, resulting in the corrosion of the anode. Fig. 14.10 (b) illustrates a simple galvanic cell, slightly different than in Fig. 14.10 (a). The anodic reaction is same as reaction (14.12), i.e., iron dissolves.
But, the electrons travel through the outer-circuit to produce the cathodic reaction:
Cu2+ + 2e̅ → Cu …(14.15)
If any of the above steps does not take place, the corrosion stops. For example, the electrons reaching the cathode, if are not consumed there, the potentials of the two metals try to become equal by polarisation to decrease, and to ultimately stop the corrosion of the anode.
Not only that the two metals must be in electrical contact, but the solution in contact has also to be an electrical (ionic) conductor. Distilled-water and pure dry-air are poor electrical conductors, and thus, are very seldom involved in galvanic corrosion. Natural water, deposits of moisture, aqueous solutions of salts, acids, or bases are good conductors.
Galvanic corrosion is common in the presence of one of these electrolytes. The electrical current flows over large distances, and even through the cathode which does not corrode. This distinguishes the electro-chemical corrosion from direct-chemical corrosion.
If instead of two dissimilar metals in electrical contact, a potential difference exists between one part of metal and another part of the same metal, for one or more reasons such as the presence of two different phases at these two places, corrosion cell forms. One of the phases becomes anode, and the other the cathode.
The corrosion of one at lower potential, i.e., the anode, occurs. Two main reactions occur, one at the anode and the other at the cathode.
Anodic reaction, i.e., the oxidation reaction as the electrons are produced by it, tries to destroy the anode to make it dissolve as an ion and supply electrons, such as for iron:
Fe → Fe2+ + 2 e̅ …(14.12)
The cathodic reaction is a reduction reaction as electrons are consumed here. A number of cathodic reactions are possible depending on the nature of electrolyte in contact, including given in Fig. 14.11, each consuming electrons without usually affecting the cathode. Reaction (14.16) is the predominant cathodic reaction in oxygen-free acidic electrolytes in which hydrogen gas evolves.
Reaction (14.13) is responsible for corrosion of metals in natural waters, i.e., O2 dissolved neutral environments which may be slightly acidic or alkaline. Such mediums normally produce insoluble solid products by the interaction of cations and anions such as-
Fe2+ + 2(OH) ̅ → Fe(OH)2 …(14.19)
the presence of oxygen oxidises it further to ferric hydroxide,
4Fe(OH)2 + O2 + 2H2O → 4Fe(OH)3 …(14.20)
but, it being unstable, finally forms hydrated ferric oxide (red rust),
Fe(OH)3 → Fe2 O3, H2O …(14.21)
Reaction (14.17) occurs if the metal has two valencies (such as ferrous and ferric) and accumulation of metal ions occurs in the electrolyte. Corrosion in acidic mine-waters occurs by this reaction.
In a given environment, one of the above cited reactions may take place in overall cathodic process, but may change with time if the environment changes. For example, hydrogen gas (reaction 14.16) may evolve initially in slightly acidic solutions, but when the hydrogen ions are exhausted, the dissolved oxygen may produce hydroxyl ions (reaction 14.13).
The products of anodic and cathodic reactions may interact (reaction 14.19), and is often decisive in controlling the rate of corrosion. The resultant product may go into solution, or a gas is evolved, does not inhibit the corrosion further. If an insoluble compound is formed, and covers the anodic metal surface, it would reduce the rate of additional corrosion. If the insoluble compound forms at a distance from the anodic area as illustrated in Fig. 14.10 (a), it has little protective value.
Rate of Corrosion:
The most common method, from the engineer’s point of view, of expressing rate of corrosion is the corrosion penetration rate (CPR) in mils per year (mpy) where 1 mil = 0.001 in., as-
mils per year (mpy) = 534 W / DAt
where, W is the weight loss in mg, D is the density of specimen g/cm3. A is the area in square inch, t is the exposure time in hours.
However using Faraday’s Law, the penetration rate is given as:
Corrosion penetration rate = K (ai / nd)
where, a is the atomic-weight of the metal, i is the current density in micro-amperes per cm2, n is the number of electrons lost, D is the density in g/cm3. K is a constant. For most applications, a corrosion penetration rate less than about 20 mpy (0.50 mm/yr) is acceptable.
The emf series, as a guide to galvanic-corrosion-tendencies is somewhat idealist. Engineering designs seldom use pure metals in standard concentration solutions. Instead commercial alloys are used in various aqueous environments. In such cases, emf series with standard potentials is less useful, because the metal, which corrodes preferentially, may differ because of non-equilibrium effects of a specific corrosive medium.
It is more useful to have a simple, qualitative ranking of alloys for their relative tendencies to be active, or noble in a common environment. Such a data in tabular form is called galvanic series (in that environment). As an example, the galvanic series in sea water is given in Table 14.3. For environments other than sea water, the galvanic series may differ significantly from the series in Table 14.3. Table 14.3 acts as a useful guide to the design engineer in predicting the relative behaviour of adjacent materials in marine applications.
A pair of materials from this table if connected electrically in sea water, the metal lower in the series will be anode. The series shows that alloy composition radically effects the tendency towards corrosion. Also, stainless steels (passive) having protective layer is placed near the noblest alloys.
Alloys also find places in such a series. Galvanic series cannot be used to predict the effect of pH. Galvanic series is different if the environment changes. Thus, several galvanic series might be needed for applications. This series also does not indicate the extent of polarisation of electrodes, and also the rate of corrosion.
Effect of Polarization on Working of a Corrosion Cell (Evan’s Diagrams):
The process of corrosion is due to the flow of positive metallic-ions from the metal into the solution, which could mean as a flow of electrical current, I (total current) due to potential E. The electrode potential of the anode and the cathode may be plotted against the current (total current) passed by the cell.
The latter could be taken as the rate of corrosion as illustrated in Fig. 14.24. These are called Evan’s diagrams. These diagrams are of greater significance for even local attacks, in which the anode and cathode surfaces are of different sizes (As here current and not the current density is plotted).
Ec is the standard potential of the cathode, and Ea of the anode. As the corrosion starts, the current, I starts flowing through the cell. As the current increases, Ec decreases and Ea increases due to polarisation to finally intersect each other at point p, where Icorr is the maximum corrosion current passed by the cell.
At the point p, the anode and the cathode are at the same potential, called the corrosion potential, Ecorr. At any other instant say at I’, XY is the potential of the cell. The increase of polarisation at either the cathode, or the anode decreases the lcorr, resulting in the decrease of rate of anodic corrosion, as the polarisation of one of the electrodes dominates the corrosion process. For example, a corrosion cell having aerated neutral solution as electrolyte, i.e., the oxygen is present dissolved in the neutral solution, corrosion occurs by the cathodic reaction,
H2O + 2 e̅ + 1/2 O2 → 2 (OH) ̅
which needs continuous supply of oxygen. If the current is high, more is the corrosion, as this reaction proceeds at a faster rate, requiring more and more of oxygen. The supply of oxygen becomes the rate controlling step. If less oxygen is available, the cathodic reaction proceeds slowly, causing concentration polarisation of the cathode.
The curve I in Fig. 14.25 (a) is for stagnant neutral solution, i.e., having little dissolved oxygen in it, and the curve 2 is for fast moving neutral solution in which oxygen supply is more. The cathode polarises more in stagnant solution than in the latter. The rate of corrosion is more in case 2 as the value of I”corr is greater as compared to I’corr if other factors are constant. This is cathodic control corrosion (more cathodic polarisation), where the corrosion-potential is close to equilibrium potential of the anode.
Hydrogen overvoltage can play a greater part to polarise the electrodes. If two corrosion cells are formed, one between iron and copper, and the other between iron and platinum with dilute acid as the electrolyte, iron becomes the anode in both the cells, because of its potential.
The cathodic reaction in an acid solution is hydrogen evolution reaction (14.16). But the hydrogen overvoltage on copper is more than that on platinum. More polarisation occurs of copper electrode as illustrated in Fig. 14.26 (a). This polarisation, i.e., the potential is decreased by the voltage drop across the surface of the metal (necessary to overcome the activation energy barrier opposing the electrode reactions there as illustrated in Fig. 14.27, where hydrogen atoms form a layer).
The result is that lcorr for iron-platinum cell is more than for iron-copper I’ccorr) cell. Thus, the corrosion of iron in contact with platinum takes place more rapidly. In general, it can be said that cathodic polarisation limits the corrosion current, to result in the decrease of corrosion rate. It shall be proper to discuss here the importance of relative areas of cathode and anode. If the cathodic area is small as compared to anodic area, and as both carry the same current.
The current density (A/M2) at the cathode becomes more (than at the anode), which causes greater polarisation of the cathode to help in decreasing, and finally to stop the corrosion process. On the other hand, the large cathode gets enough oxygen supply to cause severe corrosion of the anode, as the cathode is polarised to a lesser extent.
Corrosion Rates from Polarisation and Mixed Potential Theory:
The mixed potential theory is based on two hypotheses:
1. An electrochemical reaction can be separated into two or more partial oxidation and reduction reactions.
2. There can be no net accumulation of electric charge during an electro-chemical reaction, i.e., the total rate of oxidation must equal the total rate of reduction for an electrically isolated corroding metal sample.
A metal in contact with two or more oxidation-reduction systems is an example of mixed-electrode system. Theoretical polarisation curves are used to represent and clarify different corrosion processes. For example, if a piece of zinc is immersed in hydrochloric acid (1 molar HCI and Zn2+ – ions). At local anode (Fig. 14.28 a), zinc corrodes with the production of Zn2+ – ions (Zn → Zn2+ + 2 e̅).
The electrons flow in zinc from local anode to local cathode. At local cathode, H+ ions are reduced to H2 (2H+ + 2e̅ → H2). Location of anodes and cathodes change with time to result in evenly corroded zinc surface. It is possible to obtain a method for estimating the rate of corrosion reactions. The local zinc anode shows a positive overvoltage.
If the standard electrode potential of zinc, Eo = -0.763, and the corresponding exchange current density, io – 10-7 A/cm2, are used for the Zn/Zn2+ electrode, the approximate Tafel slopes can be drawn for this electrode as illustrated in Fig. 14.28 (b). Also, the H2/H+ electrode, (2H+ + 2e̅ → H2) can be approximated by a Tafel slope starting at Eo = 0.00 V, and the corresponding exchange current density, io = -10-10 A/cm2.
Zinc cannot remain at either of these two reversible potentials, but must have some other potential, so that the total rate of oxidation must equal the total rate of reduction, i.e., the potential of entire zinc surface must be the same, even at anodic and cathodic regions.
The anodic reaction develops a positive overvoltage, and the cathodic reaction a negative overvoltage, so that the mixed potential, EM, or Ecorr at the point of intersection of these polarisation curves, exits on zinc at the actual rate of corrosion, which is given by the corrosion current density, icorr, which is about 10-3 A/cm2 in this case. At this point, the rate of zinc dissolution is equal to the rate of hydrogen evolution, i.e., icorr represents the rate of dissolution of zinc, or of evolution of H2.
If β values and the exchange current densities are known, then the rate of corrosion of zinc can be found out from the polarisation data. Faradays law could be used to calculate the rate of corrosion. It is evident that at this stage, the external current is zero, but internal corrosion currents are flowing.
If iron is put in dilute HCI (1 molar), then its corrosion behaviour is shown in Fig. 14.28 (c). Although, the standard potential for iron is – 0.44 V, and that for zinc is – 0.763 V, i.e., the free energy change for the dissolution of iron is less than that of zinc, but the rate of corrosion of iron is greater than that of zinc when exposed to similar HCI electrolyte. Compare Fig. 14.28 (b) and (c), icorr of iron is higher than icorr of zinc. This is because the exchange current density for hydrogen evolution reaction on zinc surfaces is low.
Galvanic Couples:
As a simple example of galvanic couple, if a piece of platinum (an inert metal) is coupled to zinc (of same area as platinum) in an acid solution (air-free), then it has been seen that the corrosion of zinc is increased with vigorous evolution of hydrogen occurring on the platinum surface, but the rate of hydrogen evolution on zinc is decreased.
Fig. 14.29 (a) illustrates schematically the electrochemical characteristics of the system with three redox-systems. According to mixed-potential theory, at the steady-state, the total rate of oxidation must equal the total rate of reduction. But the determination of total rate of reduction requires addition of individual reduction currents pertaining to hydrogen evolution on zinc and on platinum.
Fig. 14.29 (a) illustrates that polarisation curves of hydrogen-evolution on zinc and of zinc-dissolution intersect to result in icorr (Zn), the corrosion rate of zinc in acid (air-free). The areas of platinum and zinc pieces are equal. The total rate of hydrogen evolution, icorr (Zn-Pt) is equal to the sum of the rates of hydrogen evolution on zinc, icorr (Zn) and hydrogen evolution on platinum surface, icorr (Pt).
As the exchange current density, i0, H+/H2 (Zn) is very low, but i0, H+/H2 (Pt) is very high, the total rate of hydrogen evolution is effectively equal to the rate of hydrogen evolution on platinum. Thus, the total reduction rate is plotted, which intersects the anodic oxidation curve for zinc to result in corrosion rate of zinc, icorr (Zn-Pt), in the galvanic couple.
Fig. 14.29 (a) illustrates that corrosion rate of zinc increases from icorr (Zn) to icorr (Zn- Pt). The mixed potential of zinc now becomes Ecorr(Zn- Pt), which intersects polarisation curves for the evolution of hydrogen on zinc and on platinum. Thus, the rate of hydrogen evolution on zinc decreases from IH2 (Zn) to iH2 (Zn-Pt). Vigorous hydrogen evolution occurs on platinum as the rate is now given by (Zn – Pt). The reason for the increase in corrosion rate of zinc when it is galvanically coupled with platinum is due to higher exchange current density for hydrogen evolution on platinum.
Fig. 14.29 (b) illustrates schematically electrochemical characteristics of a galvanic couple having both corroding metals with equal areas. Metal S is a bit more noble than metal T as Es+/S is more noble than ET+/T and icorr(s) is lesser than icorr(T). The total reduction rate curve and the total oxidation rate curve are drawn. Their point of intersection gives corrosion potential of the couple, Ecouple. The corrosion rate of metal T increases to icorr (S-T), but corrosion rate of metal S decreases to icorr (S-T).
Passivation:
If a small sample of iron (or steel) is immersed in concentrated HNO3 (70%) at room temperature, no reaction appears to occur. If a scratch is produced now on its surface, no reaction occurs then too. Now, if this acid is diluted to 1:1 by adding water, no reaction still occurs, and the iron sample remains inert as before.
However if this sample of iron is now scratched with a glass rod, a very rapid reaction of the sample occurs. If a fresh sample of iron is immersed in further diluted acid, vigorous reaction occurs even without scratching. Iron is inert in the concentrated HNO3, and remains in passive state even after transferring it to diluted HNO3 solution. But is active (the fresh piece of iron, or due to fresh surface exposed by scratching) in dilute HNO3. In passive state, the corrosion rate is very low (1:100000). The passive state is often relatively unstable.
Passivity is defined as a loss of chemical reactivity under certain environmental conditions. A metal is passive if it subsequently resists corrosion in an environment, i.e., is metastable thermodynamically (as there is large free energy decrease associated with its passage from the metallic state to the appropriate corrosion product).
Passivation is the process in which when the anode dissolves, there forms a tight, hard, adherent film having a thickness of less than 30 A° (which contains considerable water of hydration, i.e. could be oxide, or hydroxide, etc.) over the anode surface, which protects further corrosion.
This is a method of anodic control to improve the corrosion resistance of metals and alloys, but must be used with caution because of the possibility of a transition from the passive to the active state. The film acts as a barrier to the flow of current to cause the anodic polarisation, resulting in sharp decrease of the Icorr to the extent that corrosion virtually stops as illustrated in Fig. 14.30 (a).
The nature of electrolyte plays an important role apart from the anodic metal in question, in forming the passivating film. Stainless steels and aluminium form easily this passivating film in aerated water. Even iron forms, as explained above, a passivating layer of Fe2O3 in concentrated nitric acid.
The reducing atmosphere (electrolyte) destroys this passivating film to make the stainless steel active and no more stainless. The chloride-ions in an electrolyte breakdown the chromium oxide (passivating layer) film on stainless steels to cause pitting corrosion at the breaks. Artificial passivation may reduce the corrosion in some cases, such as by applying red lead, or zinc chromate in paint-primers for iron and steel to passivate them.
The passivation may be obtained with oxidising conditions either because of high oxygen activity (such as in primers), or because of induced electrical potential from a battery, or a galvanic source, as illustrated in Fig. 14.30 (a), but both give passivating film on the surface of the anode.
The illustrated passivation curve, obtained by a potentiostatic measurement, has the applied potential E plotted against log i, where i is the current density at the anode surface of a typical active-passive metal. The metal initially demonstrates behaviour similar to non-passivating metals, i.e., point A represents the equilibrium potential of the metal (anode) under given environmental conditions, and as the electrode potential becomes more positive, metal shows typical Tafel behaviour. The dissolution rate increases exponentially. This is the active region. AB corresponds to the normal polarisation behaviour of a normal corroding metal with reaction-
M → Mn+ + ne̅ …(14.34)
But at B, the formation of protective film is thermodynamically possible by such a reaction as-
M + H2O → MO + 2H + 2e̅ …(14.35)
At icrit, the rate of dissolution of anode and the rate of formation of protective film balance. Thus, for passivity to be attained, the current should exceed Icrit, to get more positive potential beyond point C. The rate of protective film formation is more so that at point D, the entire anode surface is covered with a continuous film, and the current drops to a very low value, ip, and the metal is said to be passivated.
The curve now, DE, the anodic polarisation curve becomes almost parallel to the potential axis indicating that process of dissolution of protective film is independent of potential. The range from D to E is called passive region. In this region, metal dissolution occurs at a constant rate through the passivating film by diffusion of Mn+ ions outwards to interact with O2-, or (OH) ̅ to thicken the film.
At more noble potentials beyond the end of passive region E, the dissolution rate again increases with increasing potential in the trans-passive region by one of the reactions:
2H2O → O2 + 4 H+ + 4 e̅ (in acidic solution) …(14.36)
4(OH) ̅ O2 + 2 H2O + 4 e̅ (in neutral, or basic solution) …(14.37)
The region EF is logarithmic-dependent, determined by the overvoltage of oxygen on the film surface. Electrons pass through the film in this region. The destruction of the passive film occurs at more positive potentials. A protective metal oxide may change to a higher valence oxide, or complex-ions, which do not have protective properties such as in chromium, and the state is then called trans-passivity.
In some cases, such as in aluminium, a porous film, 200-300 thick forms on the thin protective film (called anodising here). Some active-ions such as Cl ̅, Br ̅ or I ̅ if present in electrolyte, may break down the passive film, (Eb, breakdown potential) to cause pitting corrosion as illustrated by QR. A passive metal (alloy) is one which depicts a typical s-shaped dissolution curve.
When the potentiostat circuit is switched off, the potential drops immediately to be in the active direction. The potential EF at which the active state is re-established is called Flade or reactivation potential. It is a potential when pH = 0. Flade potential is slightly less than potential at point D. Lower (more negative) the Flade potential, easier it is to attain passivity, and greater is the stability of passive film. Metals like chromium (- 0.22V) and titanium (- 0.24V) have low Flade potentials, are easily passivated, while iron (+ 0.58V) has high flade potential.
Adding a more passive metal (Chromium) to less passive metal (iron) lowers Flade potential, makes passivation easy (10-15% chromium in steel).
A metal shows good passivity, if icrit is low, corrosion rate is low in passive state, has more negative passivation potential, has more wide passive range. As the temperature as well as H+ ion concentration is increased, the ‘5’-curve is slightly lowered and shifts towards right as illustrated in Fig. 14.30 (b).
An engineer should use the passive range of the anode. The current density of the cathode (or the respective potential) should be such that the cathode and anode curves cross in the passive region as illustrated in Fig. 14.30 (b). The corrosion shall be nil. If the curves cross in the active region of anode, the corrosion shall be as usual, Fig. 14.30 (c).
As is evident that current density at anode is the measure of rate of corrosion (or weight loss). The current density is effected by the effective potentials of metals, the extent of polarisation, ability to form protective films, pH of the solution, homogeneity of metal surface as well as of electrolyte, presence of other ions, rate of flow and turbulence, temperature, the type and extent of stresses.
Pourbaix, or Potential-pH Diagrams:
Pourbaix diagrams illustrate the combined effects of the potential and the pH of the aqueous solutions on the products of corrosion in metal-water systems. Just as in equilibrium diagram, such a diagram shows the nature of the stable phase as a function of electrode potential and pH of the solution. Fig. 14.31 illustrates a simplified version of Fe-H2O system, when iron corrodes in water. It is a typical of the diagrams available for almost every element and in different corroding solutions.
Here, three significantly different types of behaviour are indicated:
1. Immunity to corrosion
2. Corrosion
3. Protection due to formation of a layer of oxide, etc.
There are three types of straight lines here, each representing an equilibrium state:
1. Horizontal lines:
Horizontal lines (1), (2) and (9). The equilibrium is controlled by E and not by pH. For example, line 1 represents the significant equilibrium aspect of the reaction-
Fe2+ + 2e̅ ⟺ Fe …(14.38)
The line 1 is drawn at an arbitrary but useful value of E = – 0.62 V. This is the basis for constructing this diagram, because the corresponding rate of solution of iron must then be negligibly low at 10-6Fe2+ ions, (if iron is put in a solution of low Fe2+ , its dissolution is low). Consequently, at electrode potentials below line I, iron is immune to corrosion for practical purposes, and the region is marked immunity in diagram. For example, if iron is connected to zinc, the potential is below line 1. Thus, zinc corrodes anodically in preference to iron. This is an example of cathodic-protection.
At higher potential than, E = – 0.62 V, the concentration of Fe2+ ions increases rapidly, and therefore tendency of iron to corrosion also increases. At E0 (- 0.44 V), Fe2+-ion concentration is 1mole/litre. Line (2) represents the equilibrium,
Fe3 + e̅ ⟺ Fe2+ …(14.40)
i.e.. the ‘redox’ reaction in which ferric ion gets reduced to the ferrous state (In a redox reaction, electrode surface serves merely as a sink (or source) for electrons). The ‘redox’ potential of this electrode is given by Nernst equation-
2. Vertical Lines:
For example, line (3) represents the equilibrium:
Fe2+ + 2 H2O ⟺ Fe(OH)2 + 2 H+ …(14.42)
Here hydrogen-ions are involved, and there is no transfer of electrons. Thus, vertical lines are independent of E, i.e., there is no characteristic single potential for such a reaction as equation (14.42), but are dependent on pH of the solution. Equation (14.42) represents the hydrolysis of ferrous-ions to form the insoluble hydroxide. The hydroxide is probably a hydrated oxide of iron (FeOH2O).
Line (4) is determined only by the pH of the solution. This line refers to reaction
Fe3+ + 3H2O ⟺ Fe(OH)3 + 3H+ …(14.43)
which involves no electron transfer, but Fe(OH)3 forms, tending to prevent further corrosion of the metal.
3. Sloping Lines:
Such as lines (6), (5). These represent equilibrium involving both hydrogen-ions and electron transfer, i.e., are effected by pH as well as by potential. For example, line (6) represents-
If iron is held in a solution of pH = 11, and raised to a potential of – 0.5 V, it shall not corrode. It shall be passivated because an insoluble layer of Fe(OH)2 will form on its surface. A further increase in potential to + 0.2 V would still cause passivation by the formation of Fe(OH)3 (probably hydrated Fe2O3)
Line (5) represents:
Fe(OH)3 + 3H+ + e̅ = Fe2+ + 3H2O …(14.47)
and E = 1.08 – 0.178 pH …(14.48)
Equation (14.48) can be used to draw the line (5).
There is a second area of corrosion at the right of the diagram, i.e., highly alkaline solutions at low potentials can cause corrosion of iron by the formation of soluble bihypoferrite as for line (7).
HFeO2⁻ + 3H+ + 2 e̅ = Fe + 2H2O
which also depends on E and pH. Ions are produced, which arc stable in alkaline solutions. Such a corrosion occurs at high temperatures. It is fortunate that the rate becomes negligibly small at normal temperatures so that iron vessels could be used in industry for handling alkaline solutions. Under special conditions, caustic- embrittlement-corrosion occurs.
Pourbaix diagrams can predict the spontaneous direction of reactions, the stability and composition of corrosion products, and the effects of changes in the environment to prevent, or reduce the corrosion. For example, iron in ordinary water develops a potential of about, EM = – 0.45 V, which corresponds to position I in Fig. 14.31 within the ‘corrosion’ region of the Pourbaix diagram.
Three methods can be used to protect this iron as indicated by arrows. In the first method, cathodic protection can be used by lowering the electrode potential into region of immunity, say by connecting iron to zinc. The other two methods involve the formation of protective layers on the surface. In one such method, the corroding solution is made slightly alkaline, i.e., increase the pH of the solution.
EM remains almost the same, Fe(OH)2 forms which then produces stable compound Fe3O4 as a layer on the iron. Such a layer in general significantly decreases the corrosion rate. In another method, anodic protection can be used by raising the potential from outside by about 0.4 volt above EM, so that corroding iron moves into region of passivity.
Due to favourable kinetics, a passivating film similar to Fe2O3 (via Fe(OH)3) of about 20 A° thick forms. This film stops the dissolution of iron as Fe2+ ions. In some cases, it may not be necessary to have special anodic potential. Iron in strong nitric acid (EM is high here) forms the passivating film, to practically stop the corrosion. Steels and stainless steels, having more than 11% Cr develop passivity in mildly oxidizing solutions.
Passivating films can be broken in many ways as there are weak spots in the films. The underlying metal may not be uniform due to phase boundaries, inclusions, or simply because of different contours. The presence of chloride-ions in the solution helps in breakage. Once the film is broken locally, the large cathodic area induces rapid corrosion, and leads to pit formation.
The repair of the broken film is promoted if iron has alloying elements like chromium, nickel, and silicon, and also by the modification of corroding solution by the additions of such as chromate ions.
Pourbaix diagrams are mainly based on thermodynamic data, and not on kinetic data. Thus, they suffer from same limitations as the equilibrium diagrams. Diagrams give no information on the rate of corrosion. Passivity in such diagrams only signifies the existence of oxides, hydroxides, or other sparingly soluble substances, but does not indicate whether they are protective or not.
Such diagrams have been drawn only for a few environments. Also, the effect of alloying elements or inclusions in metals, or in solutions are not considered. Thus, pourbaix diagrams should be used with caution.
Problem:
If a piece of iron in acidic solution of pH = 6, is given a potential of about + 0.4 V, what is the corrosion behaviour according to Pourbaix diagram?
Solution:
Project the point for E – 0.4 V and pH = 6, it falls in area of passivation. Thus, such an iron gets covered with a layer of Fe(OH)3, i.e., gets passivated. The corrosion stops.
Problem:
Surface films as well as passivating films protect the metals against corrosion. Of these, which should be preferred and is stable?
Solution:
The surface film consists of a stable compound and is preferred. A passivating film is unstable, and can be broken down.