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Principles of Corrosion

Principles of Corrosion. Professor Grace Burke Director, Materials Performance Centre Room E3 Phone 64858 m.g.burke@manchester.ac.uk. With acknowledgement to Emeritus Prof. Bob Cottis. Learning Objectives…. After this module, you will be able to:

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Principles of Corrosion

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  1. Principles of Corrosion Professor Grace Burke Director, Materials Performance Centre Room E3 Phone 64858 m.g.burke@manchester.ac.uk With acknowledgement to Emeritus Prof. Bob Cottis

  2. Learning Objectives…. After this module, you will be able to: • Describe Corrosion (and its various forms) • Identify Anodic and Cathodic reactions • Describe the Corrosion Potential for a reaction • Define the Nernst Equation • Derive the thermodynamic stability of H2O • Understand pH effects • Describe Pourbaix Diagrams and their applications • Discuss Passivity

  3. Corrosion! • A major materials issue - everywhere! • From the Latin “corrodere” – to gnaw into pieces! • Definition: A chemical or electrochemical reaction between a material (generally a metal) and the environment that leads to a degradation/deterioration in the material and its properties (ASM International) • Economic Cost of Corrosion Worldwide: • 2010 - >$2.2 TRILLION Annually!(World Corrosion Organization)

  4. Corrosion: some historical examples! • Speculated that corrosion may have contributed to the downfall of Roman Empire! • Wine was stored in Pb-lined vessels by Roman heirarchy • Pb corrodes in the acetic acid-containing solutions (like red wine!) • Pb+2 in red wine would be routinely ingested by these affluent Romans, leading to Pb-poisoning/insanity!

  5. More Corrosion Effects • In the 1700’s, Pb coils were used in condensers in making brandy • Subsequently outlawed because it became recognised that Pb was causing illness! • In mid-1700’s, Pewter (Sn-up to 50% Pb) was recognised as a serious health hazard, so Pb was significantly reduced and other alloying elements were added to Sn. (no more Pb+2 into drinks)

  6. Examples of Corrosion • Bridges • Steel-reinforcements in concrete • Some Implants • Microbial corrosion/marine environments • Rusty metal/rails, beams, girders • Some cars; exhaust systems • Pipelines • Oil rigs & storage • Aircraft

  7. Forms of Corrosion • Uniform or general corrosion • Galvanic corrosion (2 dissimilar metals) • Pitting corrosion (localised corrosion on flat surface) • Crevice corrosion (geometry; occluded region) • Intergranular corrosion (preferential corrosion) • De-alloying (selective corrosion of an element) • Erosion corrosion (wear + corrosion) • Flow-Assisted Corrosion • Microbial-Induced Corrosion (MIC) -- Environmentally-Assisted Cracking (not pure corrosion)

  8. Corrosion: How to control it? • Select a corrosion-resistant material • Use a coating (paint) (barrier layer!) • Use an inhibitor (added to solution or surface) • Cathodic protection (external applied potential) • Design to avoid corrosion

  9. Corrosion of Metals in Acid • Zinc dissolves rapidly with hydrogen evolution Zn + 2HCl  ZnCl2 + H2 • Zinc is the base or active metal

  10. Corrosion of Platinum in Acid • Platinum does not react with acids • Platinum is known as a noble metal

  11. Zinc and platinum connected: current flows and H2 is evolved on platinum Zn Pt Connection of Platinum to Zinc Zinc and platinum not connected: no reaction on platinum HCl

  12. Connection of Platinum to Zinc Zn + 2HCl  ZnCl2 + H2 • But we can separate metal dissolution and hydrogen evolution Zn Zn2++ 2e- 2H+ + 2e- H2 • These are known as electrochemical reactions Reactions that involve both chemical change and the transfer of charge

  13. External Current Applied to Pt • Hydrogen evolution at one electrode 2H+ + 2e- H2or 2H2O + 2e-H2 + 2OH- A piece of metal in the solution • Oxygen evolution at the other electrode4OH- O2 + 2H2O + 4e-or 2H2O  O2 + 4H+ + 4e- • for Platinum – because it is a noble metal (for other metals (active) the reaction will be written in reverse!)

  14. Electrodes • Electrodes are pieces of metal on which an electrochemical reaction is occurring • An anode is an electrode on which an anodic or oxidation reaction is occurring • A cathode is an electrode on which a cathodic or reduction reaction is occurring

  15. Corrosion of Metals in Acid • Zinc dissolves rapidly with hydrogen evolution Zn + 2HCl  ZnCl2 + H2 • Zinc is the base or active metal • Fe dissolves with H2 evolution Fe + 2HCl  FeCl2+ H2 • Iron is the baseor active metal

  16. Corrosion of Fe in H2O • Anodic dissolution of Fe 2Fe + 3H2O  Fe2O3 + 3H2(standard chemical equation!) 2Fe  2Fe+2 + 4e- O2 + 2H2O + 2e-4OH- CATHODE ANODE OH- Fe+2 H2O O2 2H2O + 2e- H2 + 2OH- e- e- e- e-

  17. Corrosion • 2Fe + 3H2O  Fe2O3 + 3H2(standard chemical equation!) DGreaction = DGproducts – DGreactants DGreaction = (DGFe2O3+ 3DGH2) – (2DG0Fe+ 3DGH2O) Remember, there are other reactions included: Anodic reaction Cathodic reaction 2Fe  2Fe+2 + 4e- 2H2O + O2 + 4e- 4OH- • What is the driving force? Lowering the free energy of the system!

  18. Some Thermodynamics…. • Standard Gibbs Free Energy Change of the Reaction: DGo = -RT ln Kp (Kp is Equilib. Constant at 1 atm.) • Also: DGo = -nFE0 • where n= number of electrons transferred • F is Faraday’s Constant (~96500 Coulombs/equivalent) • E0 is the standard electrode potential UNITS: n (equivalents/mole) F (Coulombs/equivalent) E0 (Volts) (Volts Coul)/mole = Joules/mole (DGo )

  19. Some Thermodynamics…. • Standard Gibbs Free Energy Change of the Reaction: DGo = -RT lnKp(Kp is Equilib. Constant at 1 atm.) DG= -RT lnK • Also: DGo = -nFE0 (and DG = -nFE) • DGo = -RT lnKp= - nFE0

  20. Some Thermodynamics…. • Standard Gibbs Free Energy Change of the Reaction: DGo = -RT ln Kp (Kp is Equilib. Constant at 1 atm.) • Chemical Potential, m, denotes the change of DG of a substance (as a f (composition of a given species) at constant T and P and # of moles of other reactants m1 = mo1 + RT ln a1 where a1 = activity of component 1 Example: Me + O2 MeO2 • So, DG = moprod – moreact1 – moreact2+ RT ln (1/aO2) (since activities of the metal and metal oxide = 1) DG = moMeO2 – moMe – moO2– RT ln aO2 DG = DGo – RT ln pO2

  21. Electrodes • Electrodes are pieces of metal on which an electrochemical reaction is occurring • An anode is an electrode on which an anodic or oxidation reaction is occurring • A cathode is an electrode on which a cathodic or reduction reaction is occurring

  22. Anodic Reactions • Examples Zn  Zn2+ + 2e- Fe  Fe2+ + 2e- Al  Al3+ + 3e- Fe2+ Fe3+ + 3e- H2 2H+ + 2e- 2H2O  O2 + 4H+ + 4e- • Oxidation reactions produce electrons LEO:Losing Electrons is Oxidation!

  23. Cathodic Reactions • Examples O2 + 2H2O + 4e- 4OH- 2H2O + 2e- H2 + 2OH- Cu2+ + 2e- Cu Fe3+ + e- Fe2+ • Reduction reactions consume electrons GER: Gaining Electrons is Reduction

  24. Anodic and Cathodic Reactions LEO – GER! Losing Electrons is Oxidation – Gaining Electrons is Reduction

  25. Effect of Potential • The voltage of the metal with respect to the solution will affect electrochemical reactions • Voltage will affect the concentration of e- and H+ • Voltage of metal with respect to solution is known as the electrochemical potential

  26. Corrosion of zinc in acid • When zinc is placed in acid the metal will start to dissolve and hydrogen will start to be liberated according to the potential of the metal • Consider the anodic zinc dissolution reactionZn  Zn2+ + 2e-

  27. Corrosion of zinc in acid Zn  Zn2+ + 2e- Corrosion Rate Electrochemical Potential Ecorr Corrosion Potential 2H+ + 2e- H2 Rate of Reaction

  28. If the potential is above the Corrosion Potential, then it will fall Corrosion of zinc in acid At the Corrosion Potential, Ecorr, we have a stable mixed equilibrium Zn  Zn2+ + 2e- Electrochemical Potential 2H+ + 2e- H2 Rate of Reaction If the potential is below the Corrosion Potential, then it will rise

  29. Ecorr icorr Current density Corrosion of zinc in acid Then the corrosion rate may be expressed as the corrosion current density, icorr Zn  Zn2+ + 2e- Electrochemical Potential As the reaction involves transfer of charge, the rate of reaction may be expressed as a current per unit area, or current density 2H+ + 2e- H2

  30. When can corrosion occur? 2H+ + 2e- H2 • The concentration of hydrogen ions will influence the rate of the reaction • As the hydrogen ion concentration is increased, so the rate of the reaction increases

  31. When can corrosion occur? • On platinum (Pt) no metal dissolution will occur, but to balance the charge a reaction, which creates electrons, must occur • One such reaction is the reverse of the hydrogen evolution reaction: H2 2H+ + 2e-

  32. Thermodynamic Equilibrium 2H+ + 2e- H2 • The potential at which this reaction occurs is known as the equilibrium potential • The concentrations of reactants controls the rates of the forward and reverse reactions, and hence the equilibrium potential

  33. Corrosion • 2Fe + 3H2O  Fe2O3 + 3H2(standard chemical equation!) DGreaction = DGproducts – DGreactants DGreaction = (DGFe2O3+ 3DGH2) – (2DG0Fe+ 3DGH2O) Remember, there are other reactions included: Anodic reaction Cathodic reaction 2Fe  2Fe+2 + 4e- 2H2O + O2 + 4e- 4OH- • What is the driving force? Lowering the free energy of the system!

  34. Some Thermodynamics…. • Standard Gibbs Free Energy Change of the Reaction: DGo = -RT ln Kp (Kp is Equilib. Constant at 1 atm= activities of the products/activities of the reactants) • Also: DGo = -nFE0 • where n= number of electrons transferred • F is Faraday’s Constant (~96500 Coulombs/equivalent) • E0 is the standard electrode potential UNITS: n (equivalents/mole) F (Coulombs/equivalent) E0 (Volts) (Volts Coul)/mole = Joules/mole (DGo )

  35. Some Thermodynamics…. • Standard Gibbs Free Energy Change of the Reaction: DGo = -RT lnKp(Kp is Equilib. Constant at 1 atm.) DG= -RT lnK • Also: DGo = -nFE0 DG = -nFE • DGo= -RT lnKp= - nFE0 DG= DGo- RT lnK/Kp

  36. Some Thermodynamics…. • Chemical Potential, m, denotes the change of DG of a substance (as a f(composition of a given species) at constant T and P and # of moles of other reactants m1 = mo1 + RT ln a1 where a1 = activity of component 1 Example: Consider Oxidation of Metal Me: Me + O2 MeO2 • So, DG = moprod – moreact1 – moreact2+ RT ln (1/aO2) (since activities of the metal and metal oxide = 1) DG = moMeO2 – moMe – moO2– RT ln aO2 (aO2= pO2) DG = DGo – RT ln pO2

  37. Nernst Equation (1) E = Eo -(RT/nF)ln(aprod/areact) Eo = True thermodynamic equilibrium potential, NOT the free corrosion potential, Ecorr Eo obtained by calculation; standard conditions: 1 atm, and 25C (298K). The Nernst Equation relates the effective concentration (activities) of the components of the cell reaction to the standard cell potential.

  38. Nernst Equation (2) H2 = 2H+ + 2e- • The Nernst equation gives • For 1 atm. hydrogen gas F is Faraday’s Constant: 96487 coul/mol R is the Gas Constant: 8.314 Joule/K/mole

  39. Nernst Equation (3) H2 = 2H+ + 2e- So, substituting for R and F, and T=298K (25°C) E = Eo + (RT/nF) ln[(aH+)2/(aH2)] n=2; aH2 = 1 = Eo + (RT/2F) 2ln [H+] = Eo + (RT/F) ln [H+] = Eo + (RT/F) 2.303 log [H+] pH = -log[H+] = Eo - 0.0591 pH

  40. Thermodynamic Stability of H2O From Marcel Pourbaix: (using cal, K, R=1.985 cal/K) Dissociation of Water: H2O = H+ + OH- log [H+] x [OH-] = log K (which is the equilib. const) [H2O] Also: Log K = - S(activities or fugacities) 2.303RT = - S (activities or fugacities) 1363 (T=298K)

  41. Thermodynamic Stability of H2O Dissociation of Water: H2O = H+ + OH- log([H+] x [OH-]} = log K (which is the equilib. const) At 25°C: log K = (m0H+ + m0OH- )- (m0H2O) = 0 - 37595 + 56690= -14 1363 1363 [H+] x [OH-] = 10-14 log[H+] + log[OH-] = -14.00 Remember, pH = -log[H+] So, for pH = 7.00, [H+] = [OH-] pH < 7.00, [H+] > [OH-] pH > 7.00, [H+] < [OH-]

  42. More about H2O • Reduction: 2H+ + 2e- H2 • Oxidation: 2H2O  O2 + 4H+ + 4e- E00red= 2m0H+ - m0H2 = 1 = 0.000V 23060 x 2 46120 E00oxid= m0O2 + 4m0H+ - 2m0H2O = +113380 = +1.228V 23060 x 4 92240 H2 = 2H+ + 2e-: E00red = 0.000 – 0.0591 pH – 0.0295 log pH2 For pH2 = 1 atm, E0 red = 0.000 – 0.0591 pH 2H2O = O2 + 4H+ + 4e-: E0oxid = +1.228 – 0.0591 pH + 0.0147 log pO2 For pO2 = 1 atm, E0 oxid = +1.228 – 0.0591 pH

  43. Marcel Pourbaix (1904-1998) • Belgian Chemist (born in Russia while his father was working there as an engineer) • Professor at the University of Brussels • Derived Potential-pH diagrams, known as “Pourbaix Diagrams” • Constructed using the Nernst Equation • Provide a visual depiction of the relationship between thermodynamically possible phases of a system, bounded by lines representing the reactions between phases • Based on thermodynamics!

  44. H2O is stable H2 is stable The Pourbaix (E-pH) Diagram 2H2O = O2 + 4H+ + 4e- Equilibrium potential falls as pH increases 2.0 1.6 O2 is stable 1.2 O2 + 4H+ + 4e- = 2H2O 2H+ + 2e- = H2 Equilibrium potential falls as pH increases 0.8 Potential 0.4 0.0 2H+ + 2e- = H2 -0.4 -0.8 -1.2 -1.6 0 7 14 H+ OH- pH = - log [H+] pH

  45. Zn2+ stable in solution Zn metal stable Pourbaix Diagram for Zinc Equilibrium for Zn(OH)2 + 2OH-  ZnO22- + 2H2O 2.0 1.6 Equilibrium for Zn2+ + 2OH- Zn(OH)2 1.2 O2 + 4H+ + 4e- = 2H2O 0.8 ZnO22- stable in solution Zn(OH)2 stable solid Potential 0.4 0.0 2H+ + 2e- = H2 -0.4 Equilibrium for Zn + 2OH- Zn(OH)2 + 2e- Equilibrium for Zn  Zn2+ + 2e- -0.8 -1.2 -1.6 0 7 14 Equilibrium for Zn + 4OH-  ZnO22- + 2H2O + 2e- pH

  46. 2.0 1.6 1.2 0.8 Potential Passivity 0.4 0.0 Zn2+ stable in solution -0.4 -0.8 -1.2 Immunity Zn metal stable -1.6 0 7 14 Pourbaix Diagram for Zinc Corrosion O2 + 4H+ + 4e- = 2H2O Corrosion Zn(OH)2 stable solid ZnO22- stable in solution 2H+ + 2e- = H2 Corrosion is thermo-dynamically impossible pH

  47. What is Passivity? • Formation of an oxide/compound on a metal or alloy surface that is stable in the electrolyte, so that the metal is rendered “passive” in the environment (i.e., the material is “passivated”) • Passivation limits corrosion • Generally, strong oxidising conditions are required for passivation • Note: Passive films have some electrical conduction

  48. More on Passivity... • Many alloys exhibit passivity, such as Fe-Cr alloys • Fe-Cr alloys exhibit an increasing tendency to passivate as the Cr content increases! • The critical current density required for passivity in deaerated neutral solutions decreases as the Cr content is increased to 12 wt.%, beyond which it is constant • Fe-12+%Cr alloys and stainless steels are self-passivating (no externally applied current nor strongly oxidizing conditions are required for passivation)

  49. Breakdown of Passivity • Cl- ions can break down an existing passive film • It is generally difficult/impossible to passivate an alloy or metal in the presence of aggressive ions such as Cl-. • Aggressive ions affect the critical current density for passivity and may become incorporated into the oxide (and cause defects) • Local breakdown of a passive film can lead to.. Pitting Corrosion!

  50. Immunity versus Passivity • Immunity means that the metal or alloy will not corrode – Period! It is thermodynamically impossible. • Passivity means that the alloy will not corrode under specific conditions, but it is not permanent. It is still thermodynamically possible for the alloy to corrode.

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