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CORROSION PowerPoint Presentation

CORROSION

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CORROSION

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  1. CORROSION • OXIDATION • CORROSION • PREVENTION AGAINST CORROSION Principles and Prevention of Corrosion D.A. Jones Prentice-Hall, Englewood-Cliffs (1996)

  2. Attack of Environment on Materials • Metals get oxidized • Polymers react with oxygen and degrade • Ceramic refractories may dissolved in contact with molten materials • Materials may undergo irradiation damage

  3. Oxidation • Oxide is the more stable than the metal (for most metals) • Oxidation rate becomes significant usually only at high temperatures • The nature of the oxide determines the rate of oxidation

  4. For good oxidation resistance the oxide should be adherent to the surface • Adherence of the oxide = f(the volume of the oxide formed : the volume of metal consumed in the oxidation) = f(Pilling-Bedworth ratio) • PB < 1  tensile stresses in oxide film  brittle oxide cracks • PB > 1  compressive stresses in oxide film  uniformly cover metal surface and is protective • PB >> 1  too much compressive stresses in oxide film  oxide cracks

  5. If the metal is subjected to alternate heating and cooling cycles  the relative thermal expansion of the oxide vs metal determines the stability of the oxide layer • Oxides are prone to thermal spalling and can crack on rapid heating or cooling • If the oxide layer is volatile (e.g. Mo and W at high temperatures)  no protection

  6. Oxide Metal Progress of oxidation after forming the oxide layer: diffusion controlled activation energy for oxidation is activation energy for diffusion through the oxide layer Oxygen anions Metal Cations Oxidation occurs at air-oxide interface Oxidation occurs at metal-oxide interface • Diffusivity = f(nature of the oxide layer, defect structure of the oxide) • If PB >> 1 and reaction occurs at the M-O interface  expansion cannot be accommodated

  7. Oxidation resistant materials • As oxidation of most metals cannot be avoided the key is to form a protective oxide layer on the surface • The oxide layer should offer a high resistance to the diffusion of the species controlling the oxidation • The electrical conductivity of the oxide is a measure of the diffusivity of the ions (a stoichiometric oxide will have a low diffusivity) • Alloying the base metal can improve the oxidation resistance • E.g. the oxidation resistance of Fe can be improved by alloying with Cr, Al, Ni • Al, Ti have a protective oxide film and usually do not need any alloying

  8. Diffusion in Ionic crystals • Schottky and Frenkel defects (defects in thermal equilibrium) assist the diffusion process • If Frenkel defects dominate  the cation interstitial of the Frenkel defect carries the diffusion flux • If Schottky defects dominate  the cation vacancy carries the diffusion flux • Other defects in ionic crystals  impurities and off-stoichiometry  Cd2+ in NaCl crystal generates a cation vacancy  s diffusivity  Non-stoichiometric ZnO  Excess Zn2+   diffusivity of Zn2+  Non-stoichiometric FeO  cation vacancies   diffusivity of Fe2+ • Electrical conductivity  Diffusivity Frenkel defect Schottky defect • Cation (being smaller get displaced to interstitial voids • E.g. AgI, CaF2 • Pair of anion and cation vacancies • E.g. Alkali halides

  9. Alloying of Fe with Cr • A protective Cr2O3 layer forms on the surface of Fe (Cr2O3) = 0.001 (Fe2O3) • Upto 10 % Cr alloyed steel is used in oil refinery components • Cr > 12%  stainless steels  oxidation resistance upto 1000oC  turbine blades, furnace parts, valves for IC engines • Cr > 17%  oxidation resistance above 1000oC • 18-8 stainless steel (18%Cr, 8%Ni)  excellent corrosion resistance • Kanthal (24% Cr, 5.5%Al, 2%Co)  furnace windings (1300oC) Other oxidation resistant alloys • Nichrome (80%Ni, 20%Cr)  excellent oxidation resistance • Inconel (76%Ni, 16%Cr, 7%Fe)

  10. Corrosion THE ELECTRODE POTENTIAL • When an electrode (e.g. Fe) is immersed in a solvent (e.g. H2O) some metal ions leave the electrode and –ve charge builds up in the electrode • The solvent becomes +ve and the opposing electrical layers lead to a dynamic equilibrium wherein there is no further (net) dissolution of the electrode • The potential developed by the electrode in equilibrium is a property of the metal of electrode  the electrode potential • The electrode potential is measured with the electrode in contact with a solution containing an unit concentration of the ions of the same metal with the standard hydrogen electrode as the counter electrode (whose potential is taken to be zero) Metalions -ve +ve

  11. Standard electrode potential of metals Standard potential at 25oC Increasing propensity to dissolve

  12. Alloys used in service are complex and so are the electrolytes (difficult to define in terms of M+) (the environment provides the electrolyte • Metals and alloys are arranged in a qualitative scale which gives a measure of the tendency to corrode  The Galvanic Series Galvanic series More reactive

  13. Galvanic Cell e flow Anode Zn (0.76) Cathode Cu (+0.34) Cu2+ + 2e  Cu Reduction or 2H+ + 2e  H2 or O2 + 2H2O + 4e  4OH Zn  Zn2+ + 2e oxidation Zn will corrode at the expense of Cu

  14. Anodic/cathodic electrodes Anodic/cathodic phases at the microstructural level Differences in the concentration of the Metal ion How can galvanic cells form? Differences in the concentration of oxygen Difference in the residual stress levels

  15. Different phases (even of the same metal) can form a galvanic couple at the microstructural level (In steel Cementite is noble as compared to Ferrite) • Galvanic cell may be set up due to concentration differences of the metal ion in the electrolyte  A concentration cell Metal ion deficient  anodic Metal ion excess  cathodic • A concentration cell can form due to differences in oxygen concentration Oxygen deficient region  anodic Oxygen rich region  cathodic • A galvanic cell can form due to different residual stresses in the same metal Stressed region more active  anodic Stress free region  cathodic O2 + 2H2O + 4e  4OH

  16. Polarization • Anodic and Cathodic reactions lead to concentration differences near the electrodes • This leads to variation in cathode and anode potentials (towards each other)  Polarization Vcathode IR drop through the electrolyte Potential (V)→ Vcathode Steady state current Current (I)→

  17. Passivation • Iron dissolves in dilute nitric acid, but not in concentrated nitric acid • The concentrated acid oxidizes the surface of iron and produces a thin protective oxide layer (dilute acid is not able to do so) • ↑ potential of a metal electrode  ↑ in current density (I/A) • On current density reaching a critical value  fall in current density (then remains constant)  Passivation

  18. Prevention of Corrosion Basic goal   protect the metal  avoid localized corrosion • When possible chose a nobler metal • Avoid electrical / physical contact between metals with very different electrode potentials (avoid formation of a galvanic couple) • If dissimilar metals are in contact make sure that the anodic metal has a larger surface area / volume • In case of microstructural level galvanic couple, try to use a course microstructure (where possible) to reduce number of galvanic cells formed • Modify the base metal by alloying • Protect the surface by various means • Modify the fluid in contact with the metal Remove a cathodic reactant (e.g. water) Add inhibitors which from a protective layer • Cathodic protection Use a sacrificial anode (as a coating or in electrical contact) Use an external DC source in connection with a inert/expendable electrode