CHAPTER 8 Announcement: Please fill in Class Evaluation Survey (compulsory) http://goo.gl/forms/i8vcyteo0y Access from www.afendirojan.wordpress.com Corrosion 1 2 THE COST OF CORROSION ISSUES TO ADDRESS... How does corrosion occur? Which metals are most likely to corrode? What environmental parameters affect corrosion rate? How do we prevent or control corrosion? Corrosion: -- the destructive electrochemical attack of a material. -- Ex: Al Capone's ship, Sapona, off the coast of Bimini. Photos courtesy L.M. Maestas, Sandia National Labs. Used with permission. Cost: -- 4 to 5% of the Gross National Product (GNP)* -- in the U.S. this amounts to just over $400 billion/yr** * H.H. Uhlig and W.R. Revie, Corrosion and Corrosion Control: An Introduction to Corrosion Science and Engineering, 3rd ed., John Wiley and Sons, Inc., 1985. **Economic Report of the President (1998). 3 4 CORROSION IN A GRAPEFRUIT Introduction Cu (cathode) + reduction reactions 2H 2e H 2 (gas) O 2 4H 4e 2H 2 O Zn (anode) - Zn 2+ 2e - Acid oxidation reaction Zn Zn 2+ 2e Corrosion: Deterioration of a metal resulting from chemical attack by its environment. Rate of corrosion depends upon temperature and concentration of reactants and products. Metals have free electrons that setup electrochemical cells within their structure. Metals have tendency to go back to low energy state by corroding. Ceramics and polymers suffer corrosion by direct chemical attack. 5 6 1
metal, M Platinum metal, M Platinum 1/1/2016 Oxidation Reduction Reactions ELECTROCHEMICAL CORROSION A metal (E.g. Zn) placed in HCL undergoes corrosion. Zn + 2HCL ZnCl 2 + H 2 or Zn + 2 Zn 2+ + H 2 also Zn Zn 2+ + 2e - (oxidation half cell reaction) 2 + 2e - H 2 (Reduction half cell reaction) Oxidation reaction: Metals form ions at local anode. Reduction reaction: Metal is reduced in local charge at Local cathode. Oxidation and reduction takes place at same rate. Ex: consider the corrosion of zinc in an acid solution Two reactions are necessary: -- oxidation reaction: Zn Zn 2 2e -- reduction reaction: 2H 2e H 2 (gas) Oxidation reaction Zn Zn 2+ Acid Zinc 2e - flow of e - in the metal solution H 2 (gas) H reduction reaction + Adapted from Fig. 17.1, Callister & Rethwisch 8e. (Fig. 17.1 is from M.G. Fontana, Corrosion Engineering, 3rd ed., McGraw- Hill Book Company, 1986.) 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) 7 8 Standard electrode Half-Cell Potential of Metals STANDARD HYDROGEN ELECTRODE Oxidation/Reduction half cell potentials are compared with standard hydrogen ion half cell potential. Voltage of metal (E.g.-Zn) is directly measured against hydrogen half cell electrode. Anodic to hydrogen More tendency to corrode Examples:- Fe (-0.44), Na (-2.74) Cathodic to hydrogen Less tendency to corrode Examples:- Au (1.498), Cu (0.33) Two outcomes: -- Corrosion ne - e - M n+ ions -- Metal is the anode (-) o V 0 (relative to Pt) metal e - H 2 (gas) 2e - 25ºC 1M M n+ sol n 1M sol n -- Electrodeposition ne - e - e - M n+ ions 2e - H 2 (gas) 25ºC 1M M n+ sol n 1M sol n -- Metal is the cathode (+) o V metal 0 (relative to Pt) 9 Standard Electrode Potential Adapted from Fig. 17.2, Callister & Rethwisch 8e. 10 STANDARD EMF SERIES Macroscopic Galvanic Cells with 1M Electrolyte EMF series metal Au Cu Pb Sn Ni Co Cd Fe Cr Zn Al Mg Na K more anodic more cathodic V o metal +1.420 V +0.340-0.126-0.136-0.250-0.277-0.403-0.440-0.744-0.763-1.662-2.363-2.714-2.924 o DV = 0.153V Data based on Table 17.1, Callister 8e. Metal with smaller o V metal corrodes. Ex: Cd-Ni cell o o V Cd < V Ni Cd corrodes Adapted from Fig. 17.2, Callister & Rethwisch 8e. - Cd 1.0 M 1.0 M + 25ºC Ni Cd 2 + solution Ni 2+ solution 11 12 Two dissimilar metal electrodes immersed in solution of their own ions. Electrode that has more negative oxidation potential will be oxidized. Oxidized Zn Zn 2+, Cu 2+ Reduced Cu Half cell reactions are Zn Zn 2+ + 2e - E 0 = -0.763 V Cu Cu 2+ + 2e - E 0 = + 0.337 V Or Cu 2+ + 2e - Cu E 0 = -0.337 V (negative sign) Adding two reactions, Zn + Cu2+ zn2+ + Cu E 0 cell = -1.1V 2
Galvanic Cells With Electrolytes not 1M If the concentration of electrolyte surrounding anode is not I molar, driving force for corrosion is greater. There will be more negative emf half cell reaction M M n+ + ne - Nernst Equation: E E 0 0.0592 log n C ion E = Net efm of half cell E 0 = Standard emf of half cell N = Number of electrons transferred C ion = Molar concentration of ions. Galvanic Cells With Acid or Alkaline Electrolytes Consider iron and copper electrodes in acidic electrolyte. Since standard electrode potential of Fe to oxidize is 0.44, compared to 0.337 of copper, Fe Fe 2+ + 2e - Since there are no copper ions to reduce 2 + 2e - H 2 If electrolyte contains oxidizing agent O 2 + 4H 4+ + 4e - 2H 2 O If electrolyte is neutral or basic, O 2 + 2H 2 O + 4e - 4OH - 13 14 Microscopic Galvanic Cell Corrosion of Single Electrode Concentration of Galvanic Cells When single electrode is immersed in an electrolyte, microscopic cathodes and anodes are formed due to structural irregularities. Oxidation reaction occurs at local anode and reduction reaction at local cathode. If iron is immersed in oxygenated water, If a concentration cell is created by immersing 2 electrodes in electrolytes of different concentrations of same ion, electrode in dilute electrolyte will be anode. Example:- 2 Fe electrodes immersed in electrolytes of 0.001M and 0.01 M Fe 2+ electrolyte. E fe2+ = E 0 + 0.0296 log C ion Fe Fe 2+ + 2e - O 2 + 2H 2 O + 4e - 4OH - 2Fe + 2H 2 O + O 2 2Fe 2+ + 4OH - 2Fe(OH) 2 for 0.001M E Fe2+ = -0.44V + 0.0296 log 0.001 = -0.529 V for 0.01M E Fe2+ = -0.44V + 0.0296 log 0.01 = -0.499 V Since 0.529 V is more negative than 0.499 V, electrode in 0.001 M solution is anode and gets corroded. 15 16 EFFECT OF SOLUTION CONCENTRATION AND TEMPERATURE Oxygen Concentration Cells Ex: Cd-Ni cell with standard 1 M solutions V o Ni V o Cd 0.153 V - + Cd 25ºC Ni 1.0 M 1.0 M Cd 2 + solution Ni 2+ solution Ex: Cd-Ni cell with non-standard solutions o o RT X VNi VCd VNi VCd ln nf Y - + Cd X M Cd 2 + solution Reduce V Ni - V Cd by -- increasing X -- decreasing Y -- increasing T T Ni Y M Ni 2+ solution n = #e - per unit oxid/red reaction (= 2 here) F = Faraday's constant = 96,500 C/mol. 17 18 If two electrodes are immersed in electrolytes of different oxygen concentrations, electrode in low-oxygen content electrolyte is anode. Example: Two iron electrodes, one in low oxygen concentration water and another in high oxygen concentration water. Anode reaction : Fe Fe 2+ + 2e - Cathode reaction: O 2 + 2H 2 O + 4e - 4OH - Since cathode reaction requires O 2 and electrons, high concentration oxygen is cathode. 3
Grain Grain boundary Electrochemical cells Multiple Phase Electrochemical Cells Grain boundaries are more anodic and hence get corroded by electrochemical attack. Grain boundaries are at higher energy. Impurities migrate to grain boundaries. Solute segregation might cause grain boundaries to become more cathodic. Cartridge Brass Grain Boundary Grain boundary (anode) anode Grain boundary (cathode) In multiple alloys, one phase is more anodic than another. Corrosion rates are higher in multiphase alloys. Example: In pearlite gray cast iron, graphite flake is cathodic than surrounding pearlite matrix. Anodic pearlite corrodes Steel, in martensitic condition (single phase) after quenching from austenitic condition, has better corrosion resistance. Impurities in metals leads to precipitation of intermetallic phases and hence forms anodic and cathodic regions Figure 12.10 leading to corrosion. 19 20 Rate of Uniform Corrosion Corrosion Reaction and Polarization Faraday s equation: W ItM nf iatm nf W = weight of metal (g), corroded or electroplated in an aqueous solution in time t, seconds. I = Current flow A, i = current density A/cm 2 M = atomic mass of metal g/mol n = number of atoms/electron produced or consumed F = Faradays Constant, A = area Cm 2 Corrosion rate is expressed as weight loss per unit area per unit time or loss in depth per unit time. When a metal corrodes, the potentials of local cathode and anode are not at equilibrium. Polarization: Displacement of electrode potential from their equilibrium values to some intermediate value and creation of net current flow. Point A : equilibrium potential and current density of Zn Point B : equilibrium potential and current density of H Point C : Intermediate point Zn in acid solution 21 22 Activation and Concentration Polarization Passivation Activation polarization: Electrochemical reactions that are controlled by a slow step in a reaction sequence. There is a critical activation energy to surmount energy barrier associated with slowest step. Concentration polarization: Associate with electrochemical reaction and controlled by diffusion of ions. Example: Reduction rate of ions at surface is controlled by diffusion of ions into metal surface. Passivation is loss of chemical reactivity in presence of a environmental condition. Formation of surface layer of reaction products that inhibit further reaction. Oxide film theory: Passive film is always a diffusion barrier of reaction products. Adsorption theory: Passive metals are covered by chemisorbed films of oxygen. Examples:- Stainless steel, nickel alloys, titanium and aluminum alloys. 23 24 4
more anodic (active) more cathodic (inert) 1/1/2016 Polarization Curve Polarization curve shows how the potential of a metal varies with current density. As the electrode potential is made more positive, the metal behaves as an active metal. When potential reaches E pp (primary passive potential) current density decreases an hence the corrosion rate. Further increase in potential makes metal active again. Many metals do not behave as galvanic cells due to passive films. Galvanic series gives the cathodic, anodic relationship between the metals. In flowing seawater, Zinc is more active than aluminum. Series is determined experimentally for every corrosive environment. The Galvanic Series. 25 26 GALVANIC SERIES FORMS OF CORROSION Ranking of the reactivity of metals/alloys in seawater Platinum Gold Graphite Titanium Silver 316 Stainless Steel (passive) Nickel (passive) Copper Nickel (active) Tin Lead 316 Stainless Steel (active) Iron/Steel Aluminum Alloys Cadmium Zinc Magnesium Based on Table 17.2, Callister & Rethwisch 8e. (Source of Table 17.2 is M.G. Fontana, Corrosion Engineering, 3rd ed., McGraw- Hill Book Company, 1986.) 27 Uniform Attack Oxidation & reduction reactions occur uniformly over surfaces. Selective Leaching Preferred corrosion of one element/constituent [e.g., Zn from brass (Cu-Zn)]. Intergranular Corrosion along grain boundaries, often where precip. particles form. Fig. 17.18, Callister & Rethwisch 8e. g.b. prec. attacked zones Stress corrosion Corrosion at crack tips when a tensile stress is present. Forms of corrosion Erosion-corrosion Combined chemical attack and mechanical wear (e.g., pipe elbows). Pitting Downward propagation of small pits and holes. Fig. 17.17, Callister & Rethwisch 8e. (Fig. 17.17 from M.G. Fontana, Corrosion Engineering, 3rd ed., McGraw-Hill Book Company, 1986.) Galvanic Crevice Narrow and Dissimilar metals are confined spaces. physically joined in the Rivet holes presence of an electrolyte. The Fig. 17.15, Callister & Rethwisch 8e. (Fig. 17.15 more anodic metal is courtesy LaQue Center for Corrosion Technology, Inc.) corrodes. 28 Types of Corrosion Uniform or general attack corrosion: Reaction proceeds uniformly on the entire surface. Controlled by protective coatings, inhibitors and cathodic protection. Galvanic or two metal corrosion: Electrochemical reaction leads to corrosion of on metal. Zinc coatings on steel protects steel as zinc is anodic to steel and corrodes. Large cathode area to small anode area should be avoided. Pitting Corrosion Pitting: Localized corrosive attacks that produces holes or pits in a metal. Results in sudden unexpected failure as pits go undetected (covered by corrosion products). Pitting requires an initiation period and grows in direction of gravity. Pits initiate at structural and compositional heterogeneities. 29 30 Pitting of stainless steel 5
Growth of Pit Growth of pit involves dissolution of metal in pit maintaining high acidity at the bottom. Anodic reaction at the bottom and cathodic reaction at the metal surface. At bottom, metal chloride + water Metal hydroxide + free acid. Some metals (stainless steel) have better resistance than others (titanium). Crevice Corrosion Localized electrochemical corrosion in crevices and under shielded surfaces where stagnant solutions can exist. Occurs under valve gaskets, rivets and bolts in alloy systems like steel, titanium and copper alloys. Anode: M M + + e - Cathode:O 2 + 2H 2 O + 4e - 4OH - As the solution is stagnant, oxygen is used up and not replaced. Chloride ions migrate to crevice to balance positive charge and form metal hydroxide and free acid that causes corrosion. 31 32 Intergranular Corrosion Stress Corrosion Localized corrosion at and/or adjacent to highly reactive grain boundaries resulting in disintegration. When stainless steels are heated to or cooled through sensitizing temperature range (500-800 0 C) chromium carbide precipitate along grain boundaries. When exposed to corrosive environment, the region next to grain boundaries become anodic and corrode. Stress corrosion cracking (SCC): Cracking caused by combined effect of tensile stress and corrosive environment. Stress might be residual and applied. Only certain combination of alloy and environment causes SCC. Crack initiates at pit or other discontinuity. Crack propagates perpendicular to stress Crack growth stops if either stress or corrosive environment is removed. 33 34 Erosion Corrosion and Cavitation Damage Fretting Corrosion and Selective Leaching Erosion corrosion: Acceleration in rate of corrosion due to relative motion between corrosive fluid and surface. Pits, grooves, valleys appear on surface in direction of flow. Corrosion is due to abrasive action and removal of protective film. Cavitation damage: Caused by collapse of air bubbles or vapor filled cavities in a liquid near metal surface. Rapidly collapsing air bubbles produce very high pressure (60,000 PSI) and damage the surface. Occurs at metal surface when high velocity flow and pressure are present. Fretting corrosion: Occurs at interface between materials under load subjected to vibration and slip. Metal fragments get oxidized and act as abrasives between the surfaces. Selective leaching: Selective removal of one element of alloy by corrosion. Example: Dezincification Selective removal of zinc from copper and brasses. Weakens the alloy as single metal might not have same strength as the alloy. 35 36 6
Hydrogen Damage Load carrying capacity of a metallic component reduced due to interaction with atomic/molecular hydrogen. Happens in low carbon and alloy steels, aluminum alloys and titanium alloys. Cracking, blistering, hydride formation, reduced ductility (hydrogen embrittlement). Caused due to the diffusion of hydrogen into metal. Bakeout is a process applied to the component to diffuse the hydrogen out of the metal. Oxidation- Protective Oxide Films Oxides form on metals due to reaction with air. Degree to which oxide films form depends on following factors. Volume ratio of oxide to metal consumed after oxidation should be close to 1. Good adherence. High melting point of the film. Low oxide pressure. Coefficient of expansion equal to that of metal. High temperature plasticity. Low conductivity and diffusion coefficients of metal ions and oxygen. 37 38 Mechanisms of Oxidation Oxidation Rates Oxidation partial reaction: M M 2+ + 2e - Reduction partial reaction: ½ O 2 + 2e - O 2- Oxidation starts by lateral expansion of discrete oxide nuclei. Metal diffuses as electrons or cations across oxide films. Sometimes O 2- ions diffuse to oxide metal interface and electrons diffuse to oxide gas interface. Oxidation rate is expressed as weight gained per unit area. Linear oxidation behavior W = K L t W=weight gained per unit area KL = linear rate constant. T = time If ion diffusion is controlling the step (Eg Fe, Cu) W2 = K p t+c K p = Parabolic rate constant, C = constant Some metals follow logarithmic rate law W = K e Log(Ct + A) C, A = constants, K e = logarithmic rate constant Examples:- Al, Cu, Fe (at slightly elevated temperature) 39 40 CORROSION PREVENTION (i) CORROSION PREVENTION (ii) Materials Selection -- Use metals that are relatively unreactive in the corrosion environment -- e.g., Ni in basic solutions -- Use metals that passivate - These metals form a thin, adhering oxide layer that slows corrosion. Metal oxide Metal (e.g., Al, stainless steel) Lower the temperature (reduces rates of oxidation and reduction) Apply physical barriers -- e.g., films and coatings Add inhibitors (substances added to solution that decrease its reactivity) -- Slow oxidation/reduction reactions by removing reactants (e.g., remove O2 gas by reacting it w/an inhibitor). -- Slow oxidation reaction by attaching species to the surface. Cathodic (or sacrificial) protection -- Attach a more anodic material to the one to be protected. Adapted from Fig. 17.23, Callister & Rethwisch Galvanized Steel Zn 2+ zinc zinc 2e - 2e - 8e. steel steel pipe Using a sacrificial anode Cu wire e - Mg Mg 2+ anode Earth Adapted from Fig. 17.22(a), Callister & Rethwisch 8e. 41 e.g., zinc-coated nail e.g., Mg Anode 42 7
Corrosion Control Material Selection Metallic Metals: Use proper metal for particular environment. For reducing conditions, use nickel and copper alloys. For oxidizing conditions, use chromium based alloys. Nonmetallic Metals: Limit use of polymers in presence of strong inorganic acids. Ceramics have better corrosion resistance but are brittle. Coatings Metallic Coatings: Used to protect metal by separating from corrosive environment and serving as anode. Coating applied through electroplating or roll bonding. might have several layers. Inorganic coatings: Coating with steel and glass. Steel is coated with porcelain and lined with glass. Organic coatings: Organic polymers (paints and varnishes) are used for coatings. Serve as barrier but should be applied carefully. 43 44 Design Alteration Environment General design rules: Provide allowance for corrosion in thickness. Weld rather than rivet to avoid crevice corrosion. Avoid dissimilar metals that can cause galvanic corrosion. Avoid excessive stress and stress concentration. Avoid sharp bends in pipes to prevent erosion corrosion. Design tanks and containers for early draining. design so that parts can be easily replaced. Design heating systems so that hot spots do not occur. Lower the temperature Reduces reaction rate. Decrease velocity of fluids Reduces erosion corrosion. Removing oxygen from liquids reduces corrosion. Reducing ion concentration decreases corrosion rate. Adding inhibitors inhibitors are retarding catalysts and hence reduce corrosion. 45 46 Cathodic Protection Anodic Protection Electrons are supplied to the metal structure to be protected. Example: Fe in acid Fe Fe 2+ + 2e - 2 + 2e - H 2 Corrosion of Fe will be prevented if electrons are supplied to steel structure. Electrons can be supplied by external DC supply or galvanic coupling with more anodic metal. Externally impressed anodic currents form protective passive films on metal and alloy surfaces. Anodic currents are applied by potentiostat to protect metals that passivate. Current makes them more passive and decreases the corrosion rate. 47 48 8
SUMMARY Metallic corrosion involves electrochemical reactions -- electrons are given up by metals in an oxidation reaction -- these electrons are consumed in a reduction reaction Metals and alloys are ranked according to their corrosiveness in standard emf and galvanic series. Temperature and solution composition affect corrosion rates. Forms of corrosion are classified according to mechanism Corrosion may be prevented or controlled by: -- materials selection -- reducing the temperature -- applying physical barriers -- adding inhibitors -- cathodic protection 49 9