9/19/2018. Corrosion Thermodynamics 2-3. Course Outline. Guiding Principles. Why study thermodynamics? Guiding Principles

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1 Kwame Nkrumah University of Science & Technology, Kumasi, Ghana Week 1 Course Outline Topic Introduction: Reactivity types, corrosion definition, atmospheric corrosion, classification, effects, costs, risk management Corrosion Thermodynamics Corrosion thermodynamics: Corrosion reactions, cell requirements, free energy change, electrochemical potential, Nernst equation, Eh-pH (Pourbaix) diagrams, reference electrodes Corrosion kinetics: Electrical double layer, exchange current density, activation and mass transport control, mixed potential theory, polarization diagrams, passivity 7 Mid Semester Exams 2 Why study thermodynamics? Provides understanding of energy changes involved in electrochemical corrosion reactions Show how conditions may be adjusted to control corrosion Immunity Passivity Guiding Principles Corrosion is only a question of time rust never sleeps Corrosion of metals is electrochemical in nature o Reactions involve both charge (electron) and solvated proton (H ) transfer Fe 2H 2 O = Fe(OH) 2 2H 2e Any reaction which can be divided into two (or more) partial reactions of oxidation and reduction is termed electrochemical. Eg. Reaction of Zn in HCl 3 4 Guiding Principles Chemical reaction metal (iron) either changes into rust or dissolves o Acid-base reactions involving only solvated proton (H ) transfer (no electron transfer) Fe(OH) 2 = HFeO 2- H 2Fe(OH) 3 = Fe 2 O 3.3H 2 O Corrosion rate depends on both the metal (and its condition) and the environment. Cr 2 O 3 (SS) is dynamic layer and dissolves in reducing environment. Al 2 O 3 is static. Presence of moisture is important 5 Electrochemical Nature of Corrosion Corrosion of zinc in an acid solution Two reactions are necessary: -- oxidation reaction: Zn Zn 2 2e -- reduction reaction: 2H 2e H 2 (gas) H Oxidation reaction flow of e - Zn Zn 2 H in the metal Zinc H Acid 2e- H solution H H H 2 (gas) H reduction reaction Other reduction reactions in solutions with dissolved oxygen: -- acidic solution -- neutral or basic solution O 2 4H 4e 2H 2 O O 2 2H 2 O 4e 4(OH) 6 1

Corrosion Cell Four Prerequisites for Corrosion 1. An anode on a metal surface (oxidation rxn). e.g. Fe = Fe 2 2e - 2. A cathode on a metal surface (reduction rxn). e.g. O 2 2H 2 O 4e - = 4OH - 3.

An electrical connection between anode and cathode (allows electrons to flow between anode and cathode) Importance of the Four Prerequisites of Corrosion Corrosion Control: Stopping the anodic

2 Corrosion Cell Four Prerequisites for Corrosion 1. An anode on a metal surface (oxidation rxn). e.g. Fe = Fe 2 2e - 2. A cathode on a metal surface (reduction rxn). e.g. O 2 2H 2 O 4e - = 4OH - 3. Electrolyte in contact with anode and cathode (path for ionic conduction). 4. An electrical connection between anode and cathode (allows electrons to flow between anode and cathode) Importance of the Four Prerequisites of Corrosion Corrosion Control: Stopping the anodic reaction (e.g. passivation, cathodic protection) Stopping the cathodic reaction (e.g. removing dissolved oxygen) Stopping ion flow between anodes and cathodes (e.g. organic coatings) Why do corrosion cells form? Corrosion cells form due to the difference in energy between the metal and the environment. Variations could be Metallurgical composition, microstructure, inclusions, heat treatment, welding, fabrication etc. Environmental temperature induced corrosion, microbial induced corrosion, environmental induced SCC etc. Etching as Corrosion Impurity atom/compound 2

Multiphase structure Anodic Processes Consider similar cell reactions: Zn 2HCl ZnCl 2 H 2 Zn H 2 SO 4 ZnSO 4 H 2 Fe 2HCl FeCl 2 H 2 2Al 6HCl 2AlCl 3 3H 2 Separation into anodic reactions: Zn Zn 2 2e-

reaction) Generalized anodic reaction: M M z ze Cathodic Processes Hydrogen Evolution (Deaerated Acids): 2H 2e- H 2 Oxygen Reduction: O 2 4H 4e- H 2 O (Acidic) O 2 2H 2 O 4e- 4OH - (Neutral/Alkaline)

3 Multiphase structure Anodic Processes Consider similar cell reactions: Zn 2HCl ZnCl 2 H 2 Zn H 2 SO 4 ZnSO 4 H 2 Fe 2HCl FeCl 2 H 2 2Al 6HCl 2AlCl 3 3H 2 Separation into anodic reactions: Zn Zn 2 2e- Fe Fe 2 2e- Al Al 3 3e- Some Key Anodic Processes Dissolution-Precipitation M M z ze- (electrochemical reaction) -- M z zh 2 O = M(OH) z zh (chemical reaction) Example: Pb H 2 SO 4 PbSO 4 H 2 (cell reaction) Generalized anodic reaction: M M z ze Cathodic Processes Hydrogen Evolution (Deaerated Acids): 2H 2e- H 2 Oxygen Reduction: O 2 4H 4e- H 2 O (Acidic) O 2 2H 2 O 4e- 4OH - (Neutral/Alkaline) Metal Ion Reduction: Fe 3 e- Fe 2 Metal Deposition: Cu 2 2e- Cu Direct Film Formation M H 2 O = MO ads 2H 2e- M H 2 O = MO (oxide) 2H 2e- Example: 2Cr 3H 2 O = Cr 2 O 3 (s) 6H 6e Types of Electrochemical Cells Voltaic Cell (Galvanic Cell) 3

direction written? 1. Zn/HCl cell: Anode: Zn = Zn 2 2e; E 0 a = -E 0 c = -(-0.762V) Cathode: 2H 2e = H 2 ; E 0 c = 0V E 0 (cell) = E 0 (anode) E 0 (cathode) = 0.762V Spontaneous as written 2.

4 Electrolytic Cell Galvanic Cell vs. Electrolytic Cell Standard Cell Reaction Direction Consider two corrosion cell reactions: Zn(s) HCl(aq) = ZnCl 2 (aq) H 2 (g) Cu(s) HCl(aq) = CuCl 2 (aq) H 2 (g) Are the reactions spontaneous in the direction written? 1. Zn/HCl cell: Anode: Zn = Zn 2 2e; E 0 a = -E 0 c = -(-0.762V) Cathode: 2H 2e = H 2 ; E 0 c = 0V E 0 (cell) = E 0 (anode) E 0 (cathode) = 0.762V Spontaneous as written 2. Cu/HCl cell: Anode: Cu = Cu 2 2e; E 0 a = -E 0 c = -(0.342V) Cathode: 2H 2e = H 2 ; E 0 c = 0V E 0 (cell) = V Not spontaneous as written Cell Potential Corrosion cell is usually represented as: Zn Zn 2 Cu 2 Cu The two half-cell reactions are: 1. Zn = Zn 2 2e 2. Cu 2 2e = Cu The overall reaction is: Zn Cu 2 = Zn 2 Cu Derivation of the Nernst Equation The free energy change, ΔG, is related to the cell potential by Faraday s law: G nfe were n = number of electrons transferred in corrosion reaction F = Faraday s constant, 96,500 C/mole E = cell potential in a given state (Volts) ΔG = Joules Under standard conditions: 0 0 G nfe The Nernst Equation Consider the following reaction: A B = C D G G 0 [ C][ G G RT ln{ } [ A][ 0 [ C][ nfe nfe RT ln{ } [ A][ E E 0 0 RT ln K RT [ C][ ln{ } nf [ A][ One of the most fundamental equations in corrosion engineering! 4

5 The Nernst Equation Under standard conditions: T = 298K, R = J(mol.K) -1 Nernst equation becomes [ C][ E E log{ } n [ A][ E = non equilibrium potential EMF series metal Au Cu Pb Sn Ni Co Fe Cr Zn Al Mg Na K more anodic more cathodic Standard EMF Series V o metal V o V = 0.153V Metal with smaller Vo metal corrodes. Ex: o -Ni cell V < V o Ni corrodes - 25 C Ni 1.0 M 1.0 M 2 solution Ni 2 solution 26 Solution concentration and temperature Ex: -Ni cell with standard 1 M solutions V o Ni V o - 25 C 1.0 M 2 solution V Ni 1.0 M Ni 2 solution G = -nf E Ex: -Ni cell with non-standard solutions o o RT X VNi V VNi V ln - nf Y n = #e - per unit oxid/red T Ni reaction (= 2 here) X M 2 solution Y M Ni 2 solution F = Faraday's constant = 96,500 C/mol. 27 The Nernst Equation Question What is the thermodynamic tendency for tin (Sn) to corrode in deaerated sulphuric acid (H 2 SO 4 ) at ph = 2, activity of Sn = 10-6, p H2 = 1.0 atm, at 25 C? E 0 = V Importance of Potential Interfacial potential (E) Potential of corroding metal minus potential in electrolyte (next to metal surface) Measurement of Cell Potential V Importance Potential can be measured readily Potential affects rate of corrosion Require reference electrode for measurement M Electrolyte Ref Porous tip 5

6 Common Reference Electrodes Conversion between reference electrodes 32 Hydrogen Evolution Reaction Neutral or Alkaline Environment Standard State, e c = 0 V Half Cell Potential = F(pH) Oxygen Evolution Reaction Standard State, e c = V 6

7 Neutral or Alkaline Environment Half Cell Potential = F(pH) Stability of H 2 O 7

Electrochemical System

Electrochemical System Topic Outcomes Week Topic Topic Outcomes 8-10 Electrochemical systems It is expected that students are able to: Electrochemical system and its thermodynamics Chemical reactions in