Thermodynamics of corrosion

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Corrosion Thermodynamics

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Corrosion thermodynamics

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Corrosion ThermodynamicsModule Four of CCE 281 Corrosion: Impact, Principles, and Practical Solutions

Lesson Objectives

Examine the relation between energy and surface potentialDiscuss standard potentials in relation to the stability of chemical speciesDescribe the main reference electrodes used to measure corrosion potentialExplain the features and functions of E-pH (Pourbaix) diagramsDescribe the stability of water, iron, and aluminum from a thermodynamic point of view

Required Reading

This Module consists of seventeen Web pages of required reading. The pagination is visible at the bottom of each pagewith direct links to adjacent pages.

Additional information can be found in sections 4.1 to 4.8 of the reference textbook (Corrosion Engineering: Principlesand Practice).

Introduction

One can use thermodynamics, e.g. Pourbaix or E-pH diagrams, to evaluate the theoretical activity of a given metal oralloy in a corrosion situation provided the chemical make-up of the environment is known. But for practical situations,it is important to realize that the environment is a variable that can change with time and conditions. It is alsoimportant to realize that the environment that actually affects a metal corresponds to the micro-environmentalconditions this metal really 'sees', i.e. the local environment at the surface of the metal. (reference)

It is indeed the reactivity of this local environment that will determine the real corrosion damage. Thus, an experimentthat would investigate only the nominal environmental condition without consideration for local effects such as flow,pH cells, deposits, and galvanic effects is useless for lifetime prediction.

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Information Module

Introduction

Corrosion thermodynamics

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Free Energy of a corrosion reactionStandard electrode potentialsNernst equationReference Half-Cells (Electrodes)

Conversion between reference half-cells electrodesSilver/silver chloride reference electrodeCopper/copper sulfate reference electrode

Measuring the corrosion potentialMeasuring pH

pH glass electrodespH antimony electrode

Potential-pH DiagramAluminum E-pH (Pourbaix) DiagramE-pH diagram of waterE-pH diagram of metalsIron E-pH (Pourbaix) Diagram

See also CCE 513: Corrosion Engineering

Free energy of a corrosion reaction

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Module Four of CCE 281 Corrosion: Impact, Principles, and Practical Solutions

Free Energy of a Corrosion ReactionIn electrical and electrochemical processes, electrical work is defined as the product of charges moved (Q) times thepotential (E) through which it is moved. If this work is done in an electrochemical cell in which the potentialdifference between its two half-cells is E, and the charge is that of one mole of reaction in which n moles of electronsare transferred, then the electrical work (-w) done by the cell must be nE. In this relationship, the Faraday constant Fis required to obtain coulombs from moles of electrons. In an electrochemical cell at equilibrium, no current flows andthe energy change occurring in a reaction is expressed in equation: (reference)

Under standard conditions, the standard free energy of the cell reaction DG0 is directly related to the standard potentialdifference across the cell, E0:

For solid or liquid compounds or elements, standard conditions are the pure compound or element; for gases they are100 kPa pressure; and for solutes they are the ideal 1 molar (mol/L) concentration.

Electrode potentials can be combined algebraically to give cell potential. For a galvanic cell, such as the Daniell cell,a positive cell voltage will be obtained if the difference is taken in the usual way, as equation.

The free energy change in a galvanic cell, or in a spontaneous cell reaction, is negative and the cell voltage positive.The opposite is true in an electrolytic cell that requires the application of an external potential to drive the electrolysisreaction, in which case Ecell would be negative.

Example problem 4.1

What is the significance of a negative cell potential?

Example problem 4.2

Is it possible to use the power coming out of a half cell? Explain your answer.

Free energy of a corrosion reaction

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Other thermodynamic quantities can be derived from electrochemical measurements. For example, the entropy change(DS) in a cell reaction is given by the temperature dependence of DG:

hence

and

where DH is the enthalpy change and T the absolute temperature (K).

The equilibrium constant (Keq) for the same reaction can be obtained with the following equation:

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Standard electrode potentials

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Module Four of CCE 281 Corrosion: Impact, Principles, and Practical Solutions

Standard Electrode PotentialsThe potential difference across an electrochemical cell is the potential difference measured between two electronicconductors connected to the electrodes. In the external circuit, the electrons will flow from the most negative point tothe most positive point and, by convention, the current will flow in the opposite direction. Since the electrode potentialcan be either positive or negative, the electrons in the external circuit can also be said to flow from the least positiveelectrode to the most positive electrode. A voltmeter may be used to measure the potential differences acrosselectrochemical cells but cannot measure directly the actual potential of any single electrode. Nevertheless, it isconvenient to assign part of the cell potential to one electrode and part to the other. (reference)

There are several potential bench marks in common use, but the most ancient is the half-cell in which hydrogen gas isbubbled over a platinum electrode immersed in a solution having a known concentration of hydrogen ions.

This historically important reference electrode is called the standard hydrogen electrode (SHE) if a standard solutionof acid is used. "By definition" the equilibrium potential of this electrode is zero at any temperature." The SHE is alsocalled by many "normal hydrogen electrode" (NHE) in reference to a solution containing one equivalent of protons.Strictly speaking, one must use unit activity rather than concentration of hydrogen ions and unit fugacity rather thanunit pressure of hydrogen gas.

However, the SHE can be somewhat inconvenient to use because of the need to supply hydrogen gas. Therefore, otherreference electrodes are much preferred for practical considerations. The potential difference across a reversible cellmade up of any electrode and a SHE is called the reversible potential of that electrode, E.If this other electrode is alsobeing operated under standard conditions of pressure and concentration, then the reversible potential difference acrossthe cell is the standard electrode potential E0 of that electrode.

Tables of standard electrode potentials in either alphabetical order or by decreasing potential values can be obtained ifany one electrode, operated under standard conditions, is designated as the standard electrode or standard referenceelectrode with which other electrodes can be compared.

Example problem 4.3

Rank the following ions in order of their thermodynamic ease of plating out of a solution: Cu2+, Co2+, Fe2+,Fe3+, Na+, Pb2+, Cu+

Since an electrochemical reaction can be written either as an oxidation or a reduction causing confusion in relation tothe sign of the potential of that reaction, a convention was adopted in Stockholm in 1953 to write the standardpotential of a reaction in reference to its reduction (E0

red) as shown again in a Table listing Standard-state reductionhalf-cell potentials in either alphabetical order or by decreasing potential values .

Standard electrode potentials

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Example problem 4.4

Rank the following elements in order of their thermodynamic ease of being oxidized in solution: Hg, Al, Fe,Au, Cr, Zn, Ag, Mg

Example problem 4.5

Using standard potentials and molarity for ion concentrations calculate the open circuit potential of thefollowing electrochemical reactions (balance the equations with water related chemical species whennecessary, i.e. H+, OH- and H2O):

a. H2O2 + Ni ? H2O + Ni2+

b. H2O + Mg2+ ? H2O2 + Mg

c. Ni + PbO2 ? Pb2+ + Ni2+

d. Al3+ + OH- ? Al + O2

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Standard-state reduction half-cell potentials in alphabetical order

Standard electrode potentials

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Standard-state reduction half-cell potentials by decreasing order of potential.

Standard electrode potentials

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Nernst equation

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Module Four of CCE 281 Corrosion: Impact, Principles, and Practical Solutions

Nernst EquationThe Nernst equation was named after the German chemist Walther Nernst who established very useful relationsbetween the energy or potential of a cell to the concentrations of participating ions. This equation can be derived fromthe equation linking free energy changes to the reaction quotient (Qreaction): (reference)

where, for a generalized equation of the form:

The capital letters A, B, M and N in equation represent respectively the reactants and products of a given reactionwhile the small letters represent the coefficients required to balance the reaction.

At equilibrium, DG = 0 and Qreaction corresponds to the equilibrium constant (Keq) described earlier.

In the case of an electrochemical reaction, substitution of the relationships DG = -nFE and DG0 = -nFE0 into theexpression of a reaction free energy and division of both sides by -nF gives the Nernst equation for an electrodereaction:

Combining constants at 25oC (298.15 K) gives the simpler form of the Nernst equation for an electrode reaction at thisstandard temperature:

In this equation, the electrode potential (E) would be the actual potential difference across a cell containing thiselectrode as a half-cell and a standard hydrogen electrode as the other half-cell. Alternatively, the relationship inequation can be used to combine two Nernst equations corresponding to two half-cell reactions into the Nernstequation for a cell reaction:

Nernst equation

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Some of the species that take part these electrode reactions are pure solid compounds and pure liquid compounds. Indilute aqueous solutions, water can be treated as a pure liquid. For pure solid compounds or pure liquid compounds,activities are constant and their values are considered to be unity. The activities of gases are usually taken as theirpartial pressures and the activities (ai) of solutes such as ions are the product of the molar concentration and theactivity coefficient of each chemical species (i):

The activity coefficient (gi) in equation can be a complex function highly dependent on a multitude of variables oftendifficult to even estimate. For this reason it is usually convenient to ignore (gi) and use the concentration term [i] as anapproximation of ai.

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Examples of Nernst equation: Water stability, Reference electrodes, Limiting current

Reference Half-Cells (Electrodes)

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Module Four of CCE 281 Corrosion: Impact, Principles, and Practical Solutions

Reference Half-Cells (Electrodes)The standard hydrogen half-cell is rather awkward to use under many circumstances in which potential measurementsare to be made. The other half-cells most frequently used in corrosion studies, along with their potentials relative to thestandard hydrogen half-cell, are listed in the following Table. (reference)

Equilibrium potential of the main reference electrodes used in corrosion, at 25oC

Most reference electrodes are used with a saturated solution and an excess of the salt crystals. The extra salt dissolvesinto the half-cell solution as some of the ions diffuse out of the reference cell body through the liquid junction duringnormal use. This extra buffer of salt extends the time before the reference cell starts to drift due to the depletion ofions as predicted by Nernst equation.

Conversion between reference half-cells electrodesSilver/silver chloride reference electrode

Reference Half-Cells (Electrodes)

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Copper/copper sulfate reference electrode

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Conversion between reference electrodes

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Module Four of CCE 281 Corrosion: Impact, Principles, and Practical Solutions

Conversion between Reference ElectrodesWhen reporting electrochemical potential measurements, it is always important to indicate which reference half-cellwas used to carry out the work. This information is required to compare these measurements to similar data that couldhave been obtained using any other reference half-cells. The scheme presented in the following Figure provides agraphical representation to help visualize some of the information listed here. (reference)

Conversion between reference electrodes

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Graphical scheme to compare potentials of the most commonly used reference electrodes

We can take the case of a measurement of the potential of a steel pipe buried in the ground, using a saturated copper-copper sulfate reference electrode (CCSRE). This might show a potential of -0.700 V measured in this way. To convertthis potential to a value on the scale in which the hydrogen electrode has a potential of zero, it is necessary to add0.318 volt to the potential that was measured, making it - 0.382 volt vs. SHE.

Conversion between reference electrodes

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Example problem 4.6

What does a measured potential value of 0.8 V vs. SHE would be if the potential had been measured with asaturated silver chloride electrode? ... with a saturated copper sulfate electrode?

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Silver/silver chloride reference electrode

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Module Four of CCE 281 Corrosion: Impact, Principles, and Practical Solutions

Silver/Silver Chloride Reference ElectrodeThe silver/silver chloride reference electrode is a widely used reference electrode because it is simple, inexpensive,very stable and non-toxic. As a laboratory electrode such as described in the following Figure, it is mainly used withsaturated potassium chloride (KCl) electrolyte, but can be used with lower concentrations such as 1 M KCl and evendirectly in seawater.

Schematic of a silver/silver chloride reference electrode

Silver/silver chloride reference electrode

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As indicated here, such changes in ionic concentrations also change the reference potential. Silver chloride is slightlysoluble in strong potassium chloride solutions, so it is sometimes recommended that the potassium chloride besaturated with silver chloride to avoid stripping the silver chloride off the silver wire.(reference)

Typical laboratory electrodes use a silver wire that is coated with a thin layer of silver chloride either by electroplatingor by dipping the wire in molten silver chloride. Industrial electrodes are fabricated using the same principle usingother geometries such as planar electrodes. When the electrode is placed in a saturated potassium chloride solution itdevelops a potential of 199 mV vs. SHE. The potential of the half-cell reaction shown in equation is determined by thechloride concentration of the solution, as defined by the Nernst equation.

The silver-silver chloride reference electrode develops a potential proportional to the chloride concentration, whether itis sodium chloride, potassium chloride, ammonium chloride or some other chloride salt and remains constant as longas the chloride concentration remains constant.

Most of reference electrodes use a saturated KClsolution with an excess of KCl crystals. The extra KCldissolves into the electrolyte as the potassium andchloride ions diffuse out through the liquid junction innormal use. This extra buffer of KCl extends the timebefore the reference cell starts to drift due to thedepletion of chloride ions in the electrolyte.

The silver-silver chloride electrode simplicity offabrication and fundamental ruggedness makes it a goodcandidate for many industrial applications where theelectrochemical potential has to been measured orcontrolled. The following commercial electrode isspecifically designed to be embedded in reinforcedconcrete. (Internet reference 74)

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Copper/copper sulfate reference electrode

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Module Four of CCE 281 Corrosion: Impact, Principles, and Practical Solutions

Copper/Copper Sulfate Reference ElectrodeCopper/copper sulfate half-cells are typically favored for potential measurements of systems buried in soils. Thefollowing Figure illustrates the principle of construction of a copper/copper sulfate reference electrode (CCSRE) usedfor soil application. (reference)

Schematic of a copper/copper sulfate reference electrode

And here is a picture of a commercial CCSRE ready for field work. What is often referred to as a pipe-to-soil potentialis actually the potential measured between the pipe and the reference electrode used to make the measurement. Thesoil itself has no standard value of potential against which the potential of a pipe can be measured independently.

Copper/copper sulfate reference electrode

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Figure 4.4 – Commercial copper/copper sulfate reference electrode.

The half cell potential of a CCSRE is dependent only upon the electrochemical equilibrium established between Cuand its ions in solution as shown in equation and in its corresponding Nernst equation.

For Cu in a saturated Cu/CuSO4 solution, this equilibrium is influenced modestly by temperature and not at all byother factors except for light. Therefore, this reference electrode has a constant half cell potential, making it reliable forfield potential measurements.

A saturated CCSRE can be fabricated with a solution of copper sulfate made with 40 g of CuSO4.5H2O in 25 mL ofdistilled water. The saturated solution should contain approximately 260 g/L of CuSO4 at 22oC.

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Measuring the corrosion potential

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Module Four of CCE 281 Corrosion: Impact, Principles, and Practical Solutions

Measuring the Corrosion PotentialThe potential of a corroding metal, often termed Ecorr, is probably the single most useful variable measured incorrosion studies as well as during the corrosion monitoring of complex field situations. It is readily measured bydetermining the voltage difference between a metal immersed in a given environment and an appropriate referenceelectrode.

The following Figure illustrates an experimental technique for measuring the corrosion potential of a metal M using alaboratory cell. This is accomplished by measuring the voltage difference between the reference electrode and themetal using a high impedance voltmeter capable to accurately measure small voltages without drawing any appreciablecurrent. (reference)

Experimental set-up to measure the corrosion potential of a specimen

Note that in Figure 4.5 the reference electrode is contained in a Luggin capillary to prevent any contamination of thereference electrode by the environment or the opposite, i.e. leaking some corrosive agent in the environment being

Measuring the corrosion potential

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monitored, while making potential measurements very close to the surface of the metal being monitored.

Example problem 4.7

What is the principle of a Luggin capillary and what are the main functions of such a device?

In measuring and reporting corrosion potentials, it is necessary to indicate the magnitude of the voltage and its sign. Inthe example shown in Figure 4.5, the corrosion potential of metal M is -0.405 V. The minus sign indicates that themetal is negative with respect to the reference electrode. However, if the metal was connected to the low point (Lo)and the reference electrode to the high point (Hi) the reading would be +0.405 V. It is customary to connect thereference electrode to the low point or the instrumental ground to avoid any confusion in reporting. Nonetheless, somemanufacturers of electrochemical equipment have done the opposite.

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Measuring pH

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Module Four of CCE 281 Corrosion: Impact, Principles, and Practical Solutions

Measuring pHMeasuring pH involve either the use of pH measuring electrodes or indicators whose colors are dependent on pH. ApH meter measures the difference in potential between a reference electrode insensitive to changes in pH and anelectrode sensitive to such changes. pH indicators based on color changes are normally used in the form of pH papers.The paper is wetted with the solution being measured and the resulting color is compared with color standards todetermine the pH. (reference)

A successful pH reading is dependent upon all components of the system being operational. Problems with any one ofthe three: electrode, meter or buffer will yield poor readings.

Electrodes: A pH electrode consists of two half-cells; an indicating electrode and a reference electrode. Mostapplications today use a combination electrode with both half cells in one body. Over 90% of pH measurementproblems are related to the improper use, storage or selection of electrodes.

Meters: A pH meter is in reality a high precision and high impedance voltmeter capable of reading small millivoltchanges from the pH electrode system. The meter is seldom the source of problems for pH measurements. Modern pHmeters have temperature compensation (either automatic or manual) to correct for variations in slope caused bychanges in temperature.

Buffers: These solutions of known pH value allow the user to adjust the system to read accurate measurements.

pH glass electrodespH antimony electrode

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pH glass electrode

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Module Four of CCE 281 Corrosion: Impact, Principles, and Practical Solutions

pH Glass ElectrodeA glass electrode is a potentiometric sensor made from glass of a specific composition. All glass pH electrodes haveextremely high electric resistance from, 50 to 500 MOhm. There are different types of pH glass electrode. Some ofthem have improved characteristics for working in very alkaline or acidic medium. But almost all electrodes canoperate in the 1 to 12 pH range. (reference)

A typical pH probe is a combination electrode, which combines both the glass and reference electrodes into one body.The pH electrode is essentially a galvanic cell. The measuring part of the electrode, the glass bulb on the bottom, iscoated both inside and out with a ~10nm layer of a hydrated gel. These two layers are separated by a layer of dryglass and the potential is created by the equilibrium in H+ ions across the membrane.

pH glass electrode

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Schematic description of a typical pH glass electrode (Courtesy Cole Parmer)

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pH antimony electrode

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Module Four of CCE 281 Corrosion: Impact, Principles, and Practical Solutions

pH Antimony ElectrodeAntimony is a unique metal with the characteristic of a direct relationship between pH and its measured potential. Thepotential difference or voltage developed between antimony and a copper/copper sulfate reference electrode variesbetween approximately 0.1 volts to 0.7 volts due to variations in the pH. (reference)

The antimony electrode must be cleaned prior to use. As with any other half-cells, special cleaning procedures must beused. Antimony is very brittle and must be treated carefully. The antimony tip must be kept smooth, and there must beno rough surface or pits.

Antimony electrode

The antimony pH electrode is particularly suited for solutions containing hydrofluoric acid (HF) since the sensor has noglass wetted parts and will not degrade in such environment.

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Potential-pH diagram

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Module Four of CCE 281 Corrosion: Impact, Principles, and Practical Solutions

Potential-pH DiagramThe stability of a metal when exposed to a given environment depends on a multitude of factors that may vary greatlywith the pH and oxidizing or reducing power f that environment. One useful concept to represent the effects ofaqueous environments on metals became known as potential-pH (E-pH) diagrams, also called predominance orPourbaix diagrams, which have been adopted universally since their introduction in late 1940's.

E-pH diagrams are typically plotted for various equilibria on normal Cartesian coordinates with potential (E) as theordinate (Y axis) and pH as the abscissa (X axis). The major uses of such diagrams, which can be constructed for allmetals, are: (reference)

Predicting whether or not corrosion can occurEstimating the composition of the corrosion products formedPredicting environmental changes which will prevent or reduce corrosion attack

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Aluminum E-pH DiagramE-pH Diagram of WaterE-pH of Diagrams of MetalsIron E-pH Diagram

Aluminum E-pH (Pourbaix) diagram

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Module Four of CCE 281 Corrosion: Impact, Principles, and Practical Solutions

Aluminum E-pH (Pourbaix) DiagramThe E-pH diagram of aluminum is one of the simplest E-pH diagrams. It will be used here to demonstrate how suchdiagrams are constructed from basic principles. In the following discussion, only four species containing the aluminumelement will be considered: (reference)

two solid species (Al and Al2O3.H2O)

two ionic species (Al3+ and AlO2-)

The first equilibrium to consider examines the possible presence of either Al3+ or AlO2- expressed in equation.

Since there is no change in valence of the aluminum present in the two ionic species considered, the associatedequilibrium is independent of the potential and the expression of that equilibrium can be derived in the followingexpression for standard conditions.

where Q is expressed in equation.

Assuming that the activity of H2O is unity and that the activities of the two ionic species are equal, one can obtain asimpler expression of the equilibrium in equation based purely on the activity of H+, equation and its logarithmic form,equation.

and if G0 is expressed in Joules and the temperature is 25oC or 298 K equation is even further simplified.

Aluminum E-pH (Pourbaix) diagram

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By using the standard thermodynamic data from the literature, it is possible to calculate that the free energy of reactionis in fact equal to 120.44 kJ mol-1 when both [Al3+] and [AlO2

-] are equal. Equation then becomes equation.

This is represented, in the E-pH diagram shown below, by a dotted vertical line separating the dominant presence ofAl3+ at low pH from the dominant presence of AlO2

- at the higher end of the pH scale.

E-pH diagram showing the soluble species of aluminum in water at 25oC

The next phase for constructing the aluminum E-pH diagram is to consider all possible reactions between the fourchemical species containing aluminum retained for this exercise, i.e. Al, Al2O3.H2O, Al3+, and AlO2

-. Thesereactions are summarized in the following Table.

Possible reactions in the Al-H2O system

Aluminum E-pH (Pourbaix) diagram

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A computer program would compare all possible interactions and rank the chemical species involved in terms of theirthermodynamic stability for all conditions of pH and potential would typically carry out this work. The followingFigure illustrates the results of such computation for aluminum in the presence of water at 25oC when the activities ofall species considered were set at value unity.

Aluminum E-pH (Pourbaix) diagram

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E-pH diagram of solid species of aluminum when the soluble species are at one molar concentration (25oC)

However, an additional consideration is necessary to make such diagrams useful for corrosion situations for which thepresence of soluble species in the environment never reaches values of the order of one molar. The following Figureillustrates the results that were computed by setting the concentrations of soluble species at decreasing values of one(100), one hundredth (10-2), one in ten thousand (10-4), and one in a million (10-6). The apparent stability of the solidspecies considered gradually recedes as lower values of soluble species are used in the calculations.

Aluminum E-pH (Pourbaix) diagram

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E-pH diagram of aluminum with four concentrations of soluble species (25oC)

It is customary to use the lowest boundary (10-6) as a practical indication of the corrosion stability of a metal and itssolid products (Figure 4.12).

Aluminum E-pH (Pourbaix) diagram

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E-pH corrosion diagram of aluminum at 25oC

The usefulness of this graphical representation of thermodynamic data for corrosion studies was discussed by Pourbaixwho showed three possible states of a metallic material:

Immunity region: In the conditions of potential and pH of that region a metal is considered to be totally immune fromcorrosion attack and safe to use. Cathodic protection may be used to bring the potential of a metal closer to that regionby forcing a cathodic shift, as shown for aluminum by the domain specified in the previous Figure (-1.0 to 1.2 V vs.CCSRE)

Passive region: In such region a metal tends to become coated with an oxide or hydroxide that may form on the metaleither as a compact and adherent film practically preventing all direct contact between the metal itself and theenvironment, or as a porous deposit which only partially prevents contact between the metal and the environment;

Corrosion region: Thermodynamic calculations indicate that, in such region of an E-pH diagram, a metal is stable asan ionic (soluble) product and therefore susceptible to corrosion attack. Experience is required to find out the extentand form of the corrosion attack that may occur in the corrosion region(s) of a Pourbaix diagram.

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See also: Equilibrium reactions of iron in water, Iron corrosion products, Iron species and their thermodynamic data,Rust chemistry, Rust converters, Steel corrosion

Aluminum E-pH (Pourbaix) diagram

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E-pH (Pourbaix) diagram of water

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Module Four of CCE 281 Corrosion: Impact, Principles, and Practical Solutions

E-pH Diagram of WaterThe following example illustrates how the stability or predominance diagram of water can be constructed from itsbasic thermodynamic information. The following equationdescribes the equilibrium between hydrogen ions andhydrogen gas in an aqueous environment: (reference)

Adding sufficient OH- to both sides of reaction results in the following equation in neutral or alkaline solutions:

At higher pH than neutral, this equation is a more appropriate representation of the situation. However, since theconcentrations of [H+] and [OH-] ions are related by the dissociation constant of water, these equations can besummarized in a Nernst equation.

that becomes equation at 25oC and the hydrogen partial pressure of value unity.

This equation and its alkaline or basic form, delineate the stability of water in a reducing environment and arerepresented in a graphical form by the sloping line (a) on the Pourbaix diagram shown below.

E-pH (Pourbaix) diagram of water

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E-pH stability diagram of water at 25oC

Below the equilibrium reaction shown as line (a) in this figure, the decomposition of H2O into hydrogen is favoredwhile it is thermodynamically stable above the same line (a). As potential becomes more positive or noble, water canbe decomposed into its other constituent, oxygen, as illustrated in equations and for respectively the acidic form andneutral or basic form of the same process.

And again these equivalent equations can be used to develop a Nernst expression of the potential in standardconditions of temperature and oxygen partial pressure of value unity.

E-pH (Pourbaix) diagram of water

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The line labeled (b) in the previous Figure represents the behavior of E vs. pH for this last equation. The chemicalbehavior of water across all possible values of potential and pH is divided into three regions. In the upper region, watercan be oxidized to produce oxygen while in the lower region it can be reduced to form hydrogen gas. Water istherefore only thermodynamically stable between lines (a) and (b). It is common practice to superimpose these twolines (a) and (b) on Pourbaix diagrams to mark the water stability boundaries.

Example problem 4.8

Some fuel cells operate by oxidizing hydrogen gas on an anode while reducing oxygen from ambient air incontact with a cathode. What would be the maximum voltage produced by such a cell running on purehydrogen and air in an acidic environment? Would it be different if pure oxygen was used instead of ambientair?

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E-pH (Pourbaix) diagram of metals

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Module Four of CCE 281 Corrosion: Impact, Principles, and Practical Solutions

E-pH Diagram of MetalsBuilding a Pourbaix or E-pH diagram to represent the stability of a metal or an alloy in a given environment is not aninsurmountable task. However it could take a few hours of your precious time to produce what is commonly auniversally accepted tool to discuss the expected behavior of metals, many Thanks to the well known Belgium scientistthat gave his name to these diagrams. The process of building these diagrams should always follow the followingsteps:

1. Study background reference material on the metal/environment of choice. For the Iron-water system we foundfour acceptable references.

2. Decide on the species that will be considered. For the Iron-water system the data representing the speciesconsidered is abundantly available.

3. Decide on the target state of the species considered. For many metals and alloys there are different levels ofhydration in the scale of stability. The Iron-water system is typically described in two states of hydration, i.e.wet and dry. The addition of extraneous soluble species such as commonly present chloride and sulfate ions cangreatly complicate the thermodynamic picture.

4. Write down the equations interrelating the chemical species corresponding to the state chosen. The twenty threeequations representing the wet state of iron species in pure water are shown here.

5. Well, now that the easy part is done, one has to go over a few sleepless nights to come to the diagram presentedhere.

Fortunately a few software systems are available to compute E-pH diagrams.

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Aluminum E-pH (Pourbaix) DiagramIron E-pH (Pourbaix) Diagram

See also: Equilibrium reactions of iron in water, Pourbaix diagram of iron, Rust chemistry

References

1. Le HH, Ghali E: Interprétation des diagrammes E-pH du systeme Fe-H2O en relation avec la fragilisationcaustique des aciers. Journal of Applied Electrochemistry 1993;72-77.

2. Silverman DC: Presence of Solid Fe(OH)2 in EMF-pH Diagram for Iron. Corrosion 1982;38:453-455.3. Townsend HE: Potential-pH Diagrams at Elevated Temperature for the System Fe-H2O. Corrosion Science

E-pH (Pourbaix) diagram of metals

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1970;10:343-358.4. Biernat RJ, Robins RG: High-Temperature Potential/pH Diagrams for the Iron-Water and Iron-Water-Sulphur

Systems. Electrochimica Acta 1972;17:1261-1283. (back)