sianda11 introduction to corrosion mechanisms

8/13/2019 SIandA11 Introduction to Corrosion Mechanisms

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Surfaces, Interfaces, and their Application Introduction to corrosion

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11 Introduction to corrosion mechanisms

11.1 Introduction

Rusting is probably the most familiar, but by far not the only form of corrosion. In contrastto mechanical damage, metal corrosion is a reaction of the material with its environment,starting from the surface of the metal. The actual corrosion reactions take place in a fewnanometers thick metal/electrolyteinterface, which does not correspond to the bulk phases oneither the metallic or the electrolyte side. Furthermore, corrosion products may be presentas a thin, well-adhering oxidic surface film, which protects the underlying metal from furthercorrosion (passive film). Studies of corrosion and passivation processes are thus closelyrelated to surface analysis.

11.2 Basic concepts of corrosion11.2.1. Definitions and interpretation

The term corrosion stands for the reaction of a material with its environment, which leads to ameasurable alteration of the material (properties, behavior), and may cause functionalimpairment (damage) of a component part or the whole system. In the case of metallicmaterials, the reaction is mostly electrochemical in nature.

There are different ways of getting involved and treat questions related to corrosion. In thislecture, the focus is placed on the first point of view:

1) The corrosion specialist(scientist) things in terms of which physico-chemical reactionsoccur on the surface. Under corrosion problems he understands a need to investigate thetype of corrosion responsible for this problem and for each type, formulate the fundamentalelectrochemical and chemical reactions that take places.

! For example: uniform corrosion, pitting, crevice corrosion, galvanic corrosion, stresscorrosion cracking, hydrogen induced cracking.

2) The engineeris concerned by the appearance of the surface and possible damages. Undercorrosion, he understands the surface modification induced by the corrosion reaction and willsearch for a way to avoid this (surface protection, other materials, modification of theenvironment).

! For example: holes due to pitting, corrosion products from uniform dissolution, cracks,

oxidation, etc.

3) Usually if the corrosion is related to a large failure or a safety issue, there is anadditional involvement of the insurances or the justice. Lawyers will determine the effectivecosts of a corrosion problem and the responsibility for these damages.

! This risk is for many companies the driving force for new development in the field ofcorrosion research and protection.

For all the corrosion investigations, it is important to be aware that we are not dealing with amaterial but a system problem (Fig. 11.1), with the different parameters listed below that needto be considered as well as some interaction between them in the corrosion mechanisms

description.

- Choice of alloying element- Heat treatment- Microstructure- Surface treatment


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Figure 11.1: The different influencing factors relevant for corrosion processes

11.2.2. Corrosion as a short-circuited galvanic element

Corrosion processes on metallic materials are, apart from few exceptions, always electro-chemical processes (Redox processes). The total reaction (Fig. 11.2) can be formally split upinto two partial reactions:

a) Oxidation reaction. This is the actual corrosion process, i.e. the metal dissolution(conversion of iron atoms from the metallic into the ionic state) the oxidation reactiontakes place at the anode:

Fe"Fe2++ 2 e-

b) Reduction reaction. Due to the electro-neutrality principle, the electrons released duringthe anodic reaction must be taken up by a part of the environment adjacent to the metal,which is then reduced. This process is taking place at the cathode. If the corrosive agent isan acidic solution, protons are reduced forming hydrogen gas:

2 H++ 2 e-"H2

(gas)

In contrast, if oxygen, dissolved in (neutral or alkaline) electrolytes, interacts with the metal,oxygen is the oxidizing agent, i.e. it will be reduced:

O2+ 2 H2O + 4 e-"4 OH-

Due to the electro-neutrality (the electrons released from the iron atom need to be taken up bythe oxidizing agent), the total corrosion process is composed of at least one oxidation and onereduction process, which must take place simultaneously (Fig. 11.2).

Material

Stresses Environment- Aggressive media- Temperature

- Potential/Current-

- Design- Mechanical bulkStresses- Internal stresses

Production- Surface stresses

Tribolo


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Figure 11.2: Schematic of a corroding metal electrode. Formal breakdown into two half-cells

of a galvanic element (oxidation reaction at the anode, reduction reaction at the cathode).

An anodic (positive) current corresponds to the iron dissolution; a cathodic (negative) currentcorresponds to the reduction reaction. Since metals (iron) are electrical conductors and theelectrolyte is in general well electrolytically conductive, both the anodic and cathodic reactionconstitutes a formally short-circuited galvanic element a current I (corrosion current) isflowing:

I = !U / (Ra+ Rc+ Re)

The intensity of the corrosion current is determined by the voltage difference !U of thegalvanic element and the resistance of the anode Ra, the cathode Rcand the electrolyte Re.

Thermodynamic and kinetic basic principles of the corrosion reaction allow the prediction ofwhether a corrosion reaction is possible or not (thermodynamics) and how fast it proceeds(kinetics). Both thermodynamic and kinetic considerations have to take both the metalas well as its environment into account.

11.2.3. Cost of corrosion

The direct costs related to corrosion degradation are estimated for industrial countries to bearound 3-4% of their GDP. For the USA, where a detailed analysis has been performed(NACE: National Association of Corrosion Engineers source), this represents 276 B$/ year =3.1% of GDP.

1998 U.S. GDP

B$8,790


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Figure 11.3: Overall corrosion costs and main industrial fields contributing to these costs.

The measures to combat corrosion are of two natures

1) Non technical

Increase awareness of the large corrosion costs and potential savings Change the misconception that nothing can be done about corrosion Change policies, regulations, standards, and management practices to increase

corrosion savings

Improve education and training of staff2) Technical

Advance design practices for better corrosion management Advance life prediction and performance assessment methods Advance corrosion technology through research, development, and implementation

11.2.4. Some preliminary considerations

In relation with the measures to combat corrosion listed previously, it is necessary to stresstwo important facts:

1) Failure of a component is measurable and is usually documented correctly, butcorrosion has initiated long before and the other consequences of corrosion processesare until now not enough investigated.

One important example is the release of corrosion products that can be toxic in:

- Medical application (implant, dentistry)

- Environmental consideration (metal interaction with living organisms)


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a) b) c)

Figure 11.4: Corrosion of biomedical implants include the aspect of ionic leaching andaccumulation in tissues: a) Cobalt-Chrome-Molybdenum hip implant, b) Co, Cr, Mo release

in tissues as function of fixation, c) Histological section showing local accumulation of toxicions. (54000 time more than the blank level)

When implants are placed in-vivo (ASTM-F75-92 test, 8.5 months implantation in sheep),and if they are not completely fixed, an important amount of corrosion products will end up inthe tissues or the blood and possibly generate allergies or even cancer, Fig. 11.4. On thisfigure, it is clearly visible that significant amount of Co, Cr, Mo are released from theimplants surface (b) and that they also tend to locally accumulate (colored areas) in the tissuenear the implant (c).

2) The choice of a wrong material because of an inadequate assessment of its corrosionresistance. Examples of common heard statements after damages are evidenced: I used a stainless Steel, it is corrosion resistant Aluminum does not corrode (the surface is never getting brown ?)

Most of these mistakes rely on the fact that it is not yet well understood in the industrialcommunity that localized attack are the most critical ones and have to be analyzed seriously.Once, such an attack has started and is growing, an autocatalytic processes that is verydifficult to control and identify takes place and will lead to later failure (Fig. 11.5). Anotheraspect that is important to mention is that these attacks do not always lead to visiblecorrosion products like rust. Formation of rust is in fact by far not the most dangerous

corrosion process.


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Figure 11.5: Complexity of the localized corrosion phenomena, from the femtoamperecurrent responsible from oxide breakdown and initiation to the macroscopic failure

a) b)

Figure 11.6: Very local electrochemical measurements with a microcapillary cell (a) ondifferent part of an MnS defect in 18Cr10Ni Stainless Steel (b)Corrosion prediction (lifetime and risk analysis) as well as materials development requires adetailed understanding of the corrosion susceptibility of any defects and microstructural partof a material or coating. New very local electrochemical methods with highest current

resolution in the femtoampere range are necessary to assess oxide film stability and corrosionkinetics at small defects (Fig. 11.6). The example of the electrochemical polarizationmeasurement on the interface between a MnS inclusion and the stainless steel shows wherethe highest corrosion susceptibility is found in this material. This specific steel system willserve as example for the discussion of localized corrosion mechanism and role of defectsduring the lecture.

Corrosion processes and their characterization are complex mechanisms requiringinterdisciplinary know-how and experimental techniques ranging from surface physics tomechanical testing. Investigation of biological interactions with metallic surfaces is a growingfield of interest related to the ionic release problematic, but also to the fact that biofilm canmodify significantly the passivation and corrosion mechanisms.


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Figure 11.7: Interdisciplinary approach necessary in corrosion science

11.3 Thermodynamic aspects related to corrosion

11.3.1. Corrosion of iron in non-oxidizing acids

Thermodynamic considerations help to understand under which conditions, a corrosionreaction is possible. Even if the corrosion rates cannot be obtained yet, it is always necessaryto analyze thermodynamic stability of each specific material-electrolyte interface. Thisprocedure is especially important for new materials development and when the use of existing

classical materials is extended in more complex media. Considering first a well-documented case of the corrosion processes of iron in acidic media, the overall reaction is:

Fe + 2HCl FeCl2+ H2

Partial reactions:anodic (oxidation) cathodic (reduction)

Fe"Fe2++ 2e- 2H++ 2e-"H2

Figure 11.8 schematically shows the different steps of a corrosion reaction at an electrode:


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Figure 11.8: Schematic description of corrosion processes of an iron bar in hydrochloric acid

The anodic as well as the cathodic partial reaction take place at the same electrode. Thismeans that both partial reactions cannot be separated, thus no external voltage can bemeasured. The electrochemical cell is short-circuited. Nevertheless, a (theoretical) cellpotential can be defined, which is calculated from the potentials of the anodic Eaand cathodicEcpartial reactions.

The equilibrium potentials of a given reaction can be calculated with the aid of the standardpotentials (tabulated) and Nernsts law:

Ea= E0+ 2.3 RT/nF ln (cMe z+)

Letters stand for: Ea normal potentialE0 standard potentialc Me z+concentration of metal ions in solutionn number of transmitted electrons

Using the common logarithm, the equation can be written as:

Ea= E0+ 0.059/n * log (cMe z+)

11.3.2. Meaning and application of the Nernst law

The Nernst equation is the starting point for assessment of the thermodynamic stability of

metallic materials in contact with electrolytes. It is also an important tool for the formulation

of Pourbaix Diagrams routinely used to assess corrosion susceptibility. For this reason, it is

important to first briefly repeat some general concepts of how and why the interface

parameters such as chemical and electrochemical energies are related to a measured potential.

This section is not meant at a derivation of fundamental thermodynamic concepts but should

allow discussing their application in corrosion and also limitations and danger of the use ofthe Nernst equation.


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First a chemical equilibrium between three compounds (A,B,C) is considered:

The reaction can be described by the following change in Gibbs free energy #G resulting

from the transformation of substance A and B to form a new compound C.

The chemical energy of each individual component iis divided in interaction energy with

similar species (i0) and an interaction term with the solution (ln X).

Thermal oxidation containing only chemical interaction terms:

In relation with the field of corrosion, such chemical equilibrium can be applied to thermaloxidation processes where the electrolyte is a gas. From the thermodynamic laws one canthen derive whether a reaction can take place or not.

#G < 0 : the reaction takes place#G > 0 : the reaction does not take place

Obviously, in this case, additional thermal energy will be produced if the reaction isspontaneously occurring or need to be furnished to the system. Gibbs free energy variationcan, for example for oxide formation, be plotted as a function of temperature in the form ofEllingham diagrams, Fig. 11.9.

Following information can be obtained directly from Fig. 11.9:

- All the oxidation reactions that will occur spontaneously (negative Gibbs free energy).Titanium and Aluminum oxides are then clearly seen as being the most stable ones.

- Which reduction of oxides can take place through oxidation of a more reactiveelement. For example, alumina is more stable than most of other element when

oxidation of alloys is considered.


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Figure 11.9: Ellingham diagram of Gibbs free energy evolution for different oxides asfunction

Electrochemical potential:

The difference between a chemical equilibrium and an electrochemical interface is that in the

latter case, charged species are moving in the system and especially across an interface such

as an electrochemical double layer. For these species, an electrical energy component has to

be added to their chemical interaction potential.

Chemical energy + Energy necessary to bring a charged species (z) on the surface

The electrochemical Double layer

Upon contact of two phases with freely moving charges, a boundary with accumulation ofopposite charges on both sides is generated and this is called the electrochemical doublelayer. In general for corrosion processes, we are dealing with a solid / liquid interface. Theexcess charges create a potential gradient through this interface (the electrode potential),Fig. 11.10.


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Figure 11.10: Electrochemical double layer structure with charge separation inducing anelectrical energy component in the measured potential

The very high electronic conductivity of the metals results in very narrow charge separation

on this side of the interface. In the electrolyte, the charge distribution extends far inside the

liquid phase. There is a distinction between the well-defined Helmoltzplane (OHP) that is

resulting from the ordering of a monolayer of charged species (water dipoles, ions) and is

responsible for most of the potential variation in concentrated electrolytes. The diffuse part

(Gouy-Chapman) is resulting from charged molecules ordering in the solution and is playing

a more important role in very dilute electrolytes for examples.

In semiconductors, a diffuse layer (space charge layer) can also develop on the solid side of

the interface.

In summary: the electrical double layer induces potential differences at the solid/liquid

interface determining the thermodynamic stability and kinetic evolution of an electrochemical

reaction.

An important fact about electrochemical equilibrium is that no energy (at constant

temperature and volume) can be generated by single processes.

As a consequence, this means that the Gibbs energy will not change during the process:


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Lets now consider, what this means for a simple electrochemical equilibrium. O indicates theoxidized species, R the reduced one.

L: liquidM: metal

First, it is necessary to emphasize the difference between a corrosion and a classicalelectrochemical process and their characterization, Fig. 11.11. In a corrosion process,thermodynamic equilibrium is initially totally undefined, because of the absence of oxidizedspecies in solution. It is furthermore constantly varying during the process because of themodification of the concentration in solution and of the electrode. For this reason,electrochemists prefer to investigate charge transfer reactions between two species in solutionon a conducting inert electrode. This difference should always be kept in mind when applyingelectrochemical concepts/methods on corroding interfaces. It is often not possible to performdata interpretation to the same accuracy/precision in the field of corrosion and some cyclic

electrochemical methods are meaningless.

a) b)

Figure 11.11: Corrosion (a) versus electrochemical processes (b) on electrodes

For this reason, the equilibrium between two species in solution will first be considered froman electrochemical point of view, like for example different iron ions.

With following Gibbs energy expression


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The expression can be developed with all the single components. For each ionic speciesinvolved, there are two chemical terms: the interaction with similar species and thelogarithmic expression of the interaction with other species in solution. For the electrons e -,there is only a chemical energy term in the solid. All the species have to cross theelectrochemical double layer; the electrical term is then constituted by the potential drop

through this interface multiplied by the species charges.

With the X= C / C0 C0: 1mol/L (standard conditions)

It is then straightforward to separate electrical and chemical energy terms. n is than thedifference between the specific charges of the oxidized ZO and reduced ZR ionic species.

F(M- L)is the work necessary to transfer one Mol of electrons (1 Farad) through theMetal-Electrolyte interface.

The potential established through the solid-liquid interface E= F(M- L) is often the onlymeasurable quantity in electrochemical systems and is directly related to chemical interactioninformation. This relation between chemical and electrical energy is nothing else then theNernst equation:

Some important comments:

- Measured potentials are directly related to chemical interaction energies, so that thereis only one potential value for a given electrochemical equilibrium process. The signinversion of the potential often found in chemistry and engineering textbooks is just aconvention used to determine galvanic cell potentials between electrodes.

- Information about the chemical stability of compounds can be derived from potentialmeasurements. This fact is of practical importance especially for heterogeneous

electrodes (real materials) where local determination of chemical energy is much moredifficult than measurement of a potential.


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Reference electrode

The measured electrode potential is directly related to the chemical energy as mentionedpreviously but is difficult to assess in an absolute way. For this reason, the electrochemicalpotential scale is always measured versus a reference electrode. The universal one, is the socalled Standard hydrogen electrode (SHE)and all the other reactions are related to thispotential value in handbooks. This electrode has however a real disadvantage, because itinvolves hydrogen gas. In practice, hydrogen at a pressure of 1 bar should be bubbled near theplatinum electrode to guarantee a constant potential, Fig. 11.12. This is not very convenientand is therefore seldom used. The better alternative is to use a so called Standard Calomelelectrode (SCE) where a mercury-chloride stable salt is in equilibrium with its anionicspecies. This electrode, called of the second kind, has the advantage to be very stablebecause of the salt stability and the fact that it can be used with a saturated chloride solution.Currently, another alternative used to replace to toxic mercury is the silver-chloride electrode

working on the same principle.

a) b)

Figure 11.12: Reference electrodes: a) Universal Standard hydrogen electrode, b) Practicallyused Standard Calomel electrode

11.3.3. Electrochemical series of metals

For an electrochemical process with multiple reactions, !G is then replaced by the cellpotential U (!G = nF!U), which can be calculated from the equilibrium potentials E aund Ecof the anodic and cathodic partial reactions, respectively:


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!U = Ea Ec

The standard potentials E0 are tabulated for standard conditions (metal immersed in a solutionof its metal ions with the concentration cMe z+= 1 Mol/l):

Partial reaction E0Me/Me z+(Volt) E'Me/Me z+(Volt)with cMe z+= 10-6mol/l

Au"Au3++ 3e- + 1.50 + 1.38

Ag"Ag++ e- + 0.80 + 0.44

Cu"Cu2++ 2e- + 0.34 + 0.16

H2"2H++ 2e- 0.00

Pb"Pb2++ 2e- - 0.13 - 0.30

Ni"Ni2++ 2e- - 0.25 - 0.42

Fe"Fe2++ 2e- - 0.44 - 0.61

Zn"Zn2++ 2e- - 0.76 - 0.94Al"Al3++ 3e- - 1.66 - 1.78

Mg"Mg2++ 2e- - 2.37 - 2.54

- "Precious" metals possess standard potentials E0> 0, meaning very strong metallic bonds.The cell potential !U is positive in combination with a hydrogen electrode, i.e. nocorrosion takes place in non-oxidizing (oxygen free) acids.

- "Non-noble" metals possess standard potentials E0 < 0, meaning that electron can beremoved easily. The cell potential !U is negative in combination with a hydrogenelectrode, i.e. corrosion in non-oxidizing acids is possible.

The standard conditions cj= cj0

are usually not fulfilled in the case of corrosion processes asmentioned before, i.e. the corrosive agent contains only traces of metal ions, e.g. c Me+= 10-6mol/l. It is better for practical calculations to use the calculated value of the potential E a (seecolumn 3 in table above).

11.3.4 Electrochemical series of nonmetals

The application of Nernsts equation to cathodic partial reactions results in the normalpotentials of the partial reactions.

Example: Hydrogen electrode

H2+ 2H2O"2H3O++ 2e- (anodic notation)

For a hydrogen partial pressure pH2 = 1 and room temperature, the hydrogen electrodepotential is only dependent on the pH-value:

Example: Oxygen electrode

4OH-"O2+ 2H2O + 4e- (anodic notation)

For an oxygen partial pressure pO2 = 1 bar and room temperature, the oxygen electrode

potential is only dependent on the pH-value:


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The following table gives an overview of the normal potentials of some electrode reactions,which are important as cathodic partial reactions in corrosion processes "electrochemicalseries of nonmetals:

Partial reaction E0(Volts)

2Cl-"Cl2+ 2e- + 1.36

2Cr3+"Cr2O72-+ 7H2O + 14H++ 6e- + 1.33

2Br-"Br2+ 2e- + 1.07

NO + 2H2O"NO3-+ 4H++ 3e- + 0.96

Fe2+"Fe3++ e- + 0.77

4OH-"O2+ 2H2= + 4e- + 0.44

Cu"Cu2++ 2e- + 0.34

H2"2H++ 2e- 0.00

11.4. Potential pH diagram (Pourbaix diagram)

The thermodynamic equilibrium data can be presented in graphical form to allow easierinterpretation of the corrosion susceptibility of materials. These diagrams were firstintroduced by Marcel Pourbaix (in 1945)

Most of the metal ions react with water and as a first step with OH -to produce solid oxide /hydroxide corrosion products that can result in some cases in the formation of a veryprotecting surface layer (passivation). Using single Nernst equation turned out to be very nonintuitive. Pourbaix Diagrams present similar information about the surface stability of anelement but as a function of two very important parameters (applied potential, pH) because:

- It is always necessary to determine the corrosion behavior of the metal as a function ofthe solution composition and of a broad pH range.

- Applied potential is also an essential parameter because the surface can be easilypolarized by the presence of a cathodic reaction

Pourbaix diagrams provide information about possible equilibria between metal, solution withdissolved metal ions, and stable oxygen compounds, as a function of the pH-value and theelectrode potential.


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Figure 11.13: Pourbaix Diagram of iron in water with different types of reaction and oxideformed as function of the potential.

Different reaction types and equilibrium are included in the diagrams. In the followingsection, the three main types corresponding to the different domains (lines I III on thediagram) will be discussed. Establishing a Pourbaix Diagram requires some assumptionsabout the reactions taking place at the solid- liquid interface, and this is the reason why evenfor simple systems, there are not unique! In our case, presence of following species has beenassumed:

Solids: Fe, Fe3O4Fe2O3Ions in solution: Fe2+,Fe3+, HFeO2

-, H2O, H+, OH-

(cj=10-6mol/l)

The first type of equilibrium is corresponding to a straight dissolution-redeposition processwithout interaction with water.


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Oxidation - reduction equilibriums with participation of water will induce a pH dependence.

Pure chemical equilibriums also induce precipitation effects and are displayed in the PourbaixDiagrams.

In addition to all the equilibriums related to metal interaction with water, there are alsocathodic reactions that are displayed in the Pourbaix Diagrams. They are also very importantto consider, because the higher their potential, the more influence they will have on triggeringa corrosion reaction. We previously mentioned that a spontaneous reaction will be possible, ifE cathodic > E anodic . Considering the figure 11.14, it can be said that much more metals will besusceptible to corrosion problems in presence of oxygen (high equilibrium potential) than ofhydrogen. Furthermore, low pHs (acidic domain) are more dangerous, from a driving force

point of view than high pHs (alkaline domain).


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Figure 11.14: Pourbaix diagram and cathodic reactions. Stability domain of water anddriven force for corrosionFollowing the comment regarding the cathodic reactions; figure 11.15 shows an example ofmaterial that is not influenced by the hydrogen reduction. In the case of copper, theelectrochemical equilibrium potential for copper dissolution- redeposition is above thehydrogen line and will be influenced mainly by oxygen reduction.

Figure 11.15: Pourbaix diagram of copper: example of element stable above the reversible

potential of hydrogen resulting in different corrosion mechanisms.


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Figure 11.16: Pourbaix diagrams for different metals. Example: Aluminum corrodes in theacidic and alkaline range; it is passive in the neutral range (pH 4-8).Figure 11.16 display a number of Pourbaix Diagrams for common materials. For engineers,three domains of pH-Potential are very important to be assessed correctly:

- Where is the material immune (can stay in the metallic state)- Where can stable oxide / hydroxidesbe formed (passivation)- In which condition can corrosion attackproceed

Pourbaix Diagrams have obviously limitations:

They tell us what canhappen, not necessarily what willhappen No information on rate of reaction can be obtained Mostly plotted for pure metals and simple electrolytes, the database for more complex

electrolytes and still incomplete and prediction even for homogeneous alloys is quite

difficult.

Practical series of elements


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Figure 11.17: Comparison between electrochemical and practical series of elements. Inparticular, note the metals Al, Cr and Ni; their behavior is more noble in practice than

expected in theory.

Thermodynamic calculations show whether and in which direction a reaction canproceed. However, they do not allow conclusions to be drawn about reaction rates.

sianda11 introduction to corrosion mechanisms

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