In the name of GOD

What is Corrosion? Reaction of a metal with its environment Aqueous corrosion reaction with water (usually containing dissolved ions) High temperature oxidation reaction with oxygen at high temperature High temperature corrosion reaction with other gases

Examples of Corrosion Rusting of steel corrosion product (rust) is solid but not protective Reaction of aluminium with water corrosion product is insoluble in water, so may be protective Burning of magnesium in air high temperature oxidation

Acids and Bases An acid is a substance that produces excess hydrogen ions (H+) when dissolved in water examples are HCl, H2SO4 A base is a substance that produces excess hydroxyl ions (OH-) when dissolved in water examples are NaOH, KOH

Corrosion of Zinc in Acid Zinc dissolves with hydrogen evolution Zn + 2HCl  ZnCl2 + H2 Zinc known as a base or active metal One atom of zinc metal plus two molecules of hydrogen chloride (hydrochloric acid) reacts to form goes to one molecule of zinc chloride plus one molecule of hydrogen gas

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

Connection of Platinum to Zinc Zinc and platinum not connected, no reaction on platinum Zinc and platinum connected, current flows and hydrogen is evolved on platinum electrons Zn Pt HCl Zn  Zn2+ + 2e- metal  metal ions + electrons Zn + 2HCl  ZnCl2 + H2 metal + acid  salt + hydrogen 2H+ + 2e-  H2 hydrogen ions + electrons  hydrogen gas

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 One atom of zinc metal two electrons in the metal one zinc ion in solution Reactions that involve both chemical change and the transfer of charge

External Current Applied to Platinum in Acid Hydrogen evolved on negative electrode 2H+ + 2e-  H2 Oxygen evolved on positive electrode 2H2O  O2 + 4H+ + 4e- Pt HCl + - Acid - chemical species that produces hydrogen ions in water Overall reaction 2H2O  2H2 + O2

External Current Applied to Platinum in Alkali Hydrogen evolved on negative electrode 2H2O + 2e-  H2 + 2OH- Oxygen evolved on positive electrode 4OH-  O2 + 2H2O + 4e- The product of [H+] times [OH-] is 10-14, so in pure water both [H+] and [OH-] are 10-7. This leads to the concept of pH, which is defined as -log[H+] Hence pH = 0 is strong acid, 7 is neutral, and 14 is strong alkali Pt NaOH + - Alkali - chemical species that produces hydroxyl ions (OH-) in water Note that H+ and OH- are in equilibrium in water: H2O  H+ + OH- Overall reaction 2H2O  2H2 + O2

External Current Applied to Platinum A piece of metal in the solution Hydrogen evolution at one electrode 2H+ + 2e-  H2 (acids) or 2H2O + 2e- H2 + 2OH- (alkalis) Oxygen evolution at the other electrode 2H2O  O2 + 4H+ + 4e- (acids) or 4OH- O2 + 2H2O + 4e- (alkalis)

Faraday’s Law Charge is related to mass of material reacted in and electrochemical reaction: 2H+ + 2e-  H2 Two hydrogen ions To produce one molecule of hydrogen gas React with two electrons

Faraday’s Constant One mole of hydrogen ions (1 g) contains Avogadro’s number (6 1023) ions Hence electrons will react with each mole of hydrogen ions Charge on the electron is 1.6  10-19 C Hence one mole of ions requires 96500 C This is known as Faraday’s constant

Faraday’s Law

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

Anodic Reactions Examples Zn  Zn2+ + 2e- zinc corrosion Fe  Fe2+ + 2e- iron corrosion Al  Al3+ + 3e- aluminium corrosion Fe2+  Fe3+ + e- ferrous ion oxidation H2  2H+ + 2e- hydrogen oxidation 2H2O  O2 + 4H+ + 4e- oxygen evolution Oxidation reactions Produce electrons

Cathodic Reactions Examples O2 + 2H2O + 4e-  4OH- oxygen reduction 2H2O + 2e-  H2 + 2OH- hydrogen evolution Cu2+ + 2e-  Cu copper plating Fe3+ + e-  Fe2+ ferric ion reduction Reduction reactions Consume electrons

Metal Ion Hydrolysis Note that metal ions may react with water (a hydrolysis reaction) e.g. Al3+ + 3H2O  Al(OH)3 + 3H+ or 2Al3+ + 3H2O  Al2O3 + 6H+ Note that in an electrochemical reaction, we have the same number of each atom on each side of the equation, and the same overall charge

Effect of Potential Electrochemical reactions involve transfer of charge Hence, we expect that the voltage of the metal with respect to the solution will affect electrochemical reactions Voltage of metal with respect to solution is known as the electrochemical potential

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 reaction Zn  Zn2+ + 2e-

Corrosion of zinc in acid If the potential is above the Corrosion Potential, then it will fall due to production of electrons Then the corrosion rate may be expressed as the corrosion current density, icorr If the potential is below the Corrosion Potential, then it will rise, due to consumption of electrons Zn  Zn2+ + 2e- Rate of Reaction At the Corrosion Potential, Ecorr, we have a stable mixed equilibrium As the reaction involves transfer of charge, the rate of reaction may be expressed as a current per unit area, or current density Electrochemical Potential icorr Ecorr 2H+ + 2e-  H2 Current density Rate of Reaction

How Fast will Corrosion Occur? Corrosion kinetics Concerned with the rates of corrosion reactions Mixed potential theory: The corrosion potential will be that potential at which the sum of all anodic (positive) and cathodic (negative) currents on the electrode is zero Polarization The change in potential that is caused by the passage of a current

Types of Polarization Activation Polarization The polarization necessary for the electrochemical reaction to go at the given rate Given by Tafel’s Law: E = potential at current i Eo = potential at current io b = Tafel slope

E-log i and Evans Diagrams Plot E against log |i|, then activation polarization gives a straight line Eo and io for the cathodic reaction Anodic reaction, Tafel slope is positive Tafel slope expressed as mV per decade of current mV log (-i2) - log (-i1) Mixed equilibrium occurs when sum of all currents is zero log |current| Electrode Potential Ecorr and icorr for the corrosion reaction Cathodic reaction, Tafel slope is negative Eo and io for the anodic reaction

Concentration Polarization Additional polarization caused by drop in concentration of a reactant at the electrode surface As concentration falls, more polarization is needed to make the current flow Eventually, no more current can flow because no more reactant can reach the metal, and a limiting current is reached

Concentration Polarization Oxygen reduction is often affected by concentration polarization Rate of cathodic oxygen reduction without concentration polarization log |current density| Electrode Potential Rate of cathodic oxygen reduction with concentration polarization Limiting current density - rate of reaction limited by availability of oxygen at the metal surface

Resistance Polarization If there is a resistance between the anode and the cathode in a cell, then the current flowing through that resistance will cause a potential drop given by Ohm’s Law: V = IR This is important for paint films and for high resistance solutions

Resistance Polarization Resistance Polarization causes potential of anode and cathode to differ due to potential drop across solution, hence corrosion current is reduced log |current density| Electrode Potential

Passivation When a passive film is formed, this causes a marked drop in current density due to the resistance of the film and its effect as a barrier to diffusion This effect is seen on the anodic curve

Passivation The rate of corrosion will be critically affected by the cathodic curve When a stable passive film has formed, the current has a steady, low value - the passive current density Rapid rate of cathodic reaction leads to passivation, and low rate of corrosion log |current density| Electrode Potential Lower rate of cathodic reaction leads to activity, and high rate of corrosion Current falls as the passive film starts to form - the active-passive transition But it may also lead to low rate of corrosion? Very slow cathodic reaction leads to low rate of corrosion Active corrosion gives normal activation polarization

Polarization Curves Iron in hydrochloric acid Electrode Potential log |current density| Electrode Potential Anodic iron dissolution Cathodic hydrogen evolution

Polarization Curves Iron in sulphuric acid log |current density| Electrode Potential Anodic iron dissolution (with active-passive transition) Oxygen evolution on passive film (or transpassive corrosion as metal is oxidised to a higher oxidation state) Cathodic hydrogen evolution

Polarization Curves Iron in aerated neutral NaCl solution log |current density| Electrode Potential Cathodic oxygen reduction Anodic iron dissolution Cathodic hydrogen evolution

Effect of pH on reaction rate Consider hydrogen evolution reaction 2H+ + 2e-  H2 The concentration of hydrogen ions will influence the rate of the reaction As the hydrogen ion concentration is increased (i.e. the solution made more acid), so the rate of the reaction increases Similarly the potential will influence the reaction - the more negative the potential the faster the reaction

Effect of pH and potential on rate of hydrogen evolution Slower Potential Faster pH

Effect of pH on reaction rate On platinum no metal dissolution will occur, but to balance the charge a reaction which creates electrons must occur If the solution contains dissolved hydrogen, the reverse of the hydrogen evolution reaction can occur: H2 2H+ + 2e-

Effect of pH on reaction rate H2 2H+ + 2e- This reaction will go faster in alkaline solution (since H+ will be removed by H+ + OH-  H2O) This reaction will go faster at more positive potentials (because electrons will be removed from metal)

Effect of pH and potential on rate of hydrogen oxidation Oxidation Faster Reduction Faster Reduction Slower Rates equal Potential Oxidation Slower Electrochemical Equilibrium pH

Thermodynamic Equilibrium 2H+ + 2e-  H2 The potential at which it occurs for a given solution composition is known as the equilibrium potential. The concentrations of reactants controls the rates of the forward and reverse reactions and hence the equilibrium potential

Thermodynamic Equilibrium ∆Gº= -nFEº ∆G= -nFE Nernst Equation:

The Nernst equation 2H+ + 2e- = H2 The Nernst equation gives For 1 atm. hydrogen gas

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

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 0.8 Zn(OH)2 stable solid Potential 0.4 Equilibrium for Zn + 2OH-  Zn(OH)2 + 2e- ZnO22- stable in solution 0.0 Zn metal stable Zn2+ stable in solution Equilibrium for Zn  Zn2+ + 2e- -0.4 Equilibrium for Zn + 4OH-  ZnO22- + 2H2O + 2e- -0.8 -1.2 -1.6 7 14

Pourbaix Diagram for Zinc Corrosion possible with oxygen reduction Potential 7 14 2.0 1.6 0.8 1.2 -0.4 0.4 0.0 -1.6 -0.8 -1.2 Corrosion Corrosion requires strong oxidising agent Passivity Corrosion Corrosion is possible, but likely to be stifled by solid corrosion product Corrosion possible with hydrogen evolution Zn(OH)2 stable solid ZnO22- stable in solution Corrosion is thermodynamically impossible Zn metal stable Zn2+ stable in solution Immunity

Pourbaix Diagram for Gold Potential 7 14 2.0 1.6 0.8 1.2 -0.4 0.4 0.0 -1.6 -0.8 -1.2 C Passivity Gold can’t corrode with oxygen reduction or hydrogen evolution Immunity Gold metal stable

Pourbaix diagram for Copper Will copper corrode in acid? Potential 7 14 2.0 1.6 0.8 1.2 -0.4 0.4 0.0 -1.6 -0.8 -1.2 Will copper corrode in neutral waters? No - hydrogen evolution only occurs below the potential for copper corrosion Cu oxides stable CuO22- stable in soln. Cu2+ stable in solution Usually it will just passivate, but corrosion can occur in slightly acid solutions Cu metal stable

Pourbaix Diagram for Iron Will iron corrode in acid? Will iron corrode in neutral waters? Potential 7 14 2.0 1.6 0.8 1.2 -0.4 0.4 0.0 -1.6 -0.8 -1.2 Will iron corrode in alkaline solution? Yes - there is a reasonably wide range of potentials where hydrogen can be evolved and iron dissolved Fe3+ Fe oxides stable Yes - although iron can form an oxide in neutral solution, it tends not to form directly on the metal, as the potential is too low, therefore it is not protective. Fe2+ stable No - iron forms a solid oxide at all potentials, and will passivate Fe metal stable

Pourbaix diagram for Aluminium Potential 7 14 1.2 0.8 0.0 0.4 -1.2 -0.4 -0.8 -2.4 -1.6 -2.0 Al3+ Al2O3 AlO2- Al

Limitations of Pourbaix Diagrams Tell us what can happen, not necessarily what will happen No information on rate of reaction Can only be plotted for pure metals and simple solutions, not for alloys

Thanks for your attention