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Corrosion Part 1

Corrosion Part 1

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Corrosion Part 1

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  1. Suez Canal University Corrosion Part 1 Dr. Eng. Hamid A. Nagy WELD-INSPECTA CO.

  2. Corrosion

  3. Driving force • Every Process to take place, we should have some driving force. • The driving force depends on the energy of the first state and that of the final state.

  4. Driving force A Barrier Energy Driving Force B

  5. Thermodynamics • Every material has two sources of its energy • Heat content, Enthalpy • Content depending on its randomness, Entropy. • We can not measure this energy directly. • So we have a reference zero value which is the hydrogen molecule formation.

  6. Thermodynamics • Now consider the transfer of a metal from state A to state B. • This can only be done if the state A has higher energy than state B. DGA > DGB DG is the free energy of material state.

  7. Thermodynamics • Now consider the reaction between two materials a and B to produce C. • The same law applies. DGA + DGB > DGC For the reaction to proceed in the direction A + B → C And vice versa.

  8. Effect of Concentration • Consider the following Reaction M++ + 2e → M • As the M ions concentration increases the reaction tends to go the left. • Oxidation reaction increases. • More anodic tendency. • Potential decreases.

  9. Kinetics • Rate of reaction depends on what is called mechanism. • There should be some energy done to activate the first stage. • This is called the energy barrier. • This energy barrier could be high or low depending on the mechanism.

  10. Kinetics • Overcoming energy barrier may consist of several steps. • The rate of occurrence of this reaction depends on the interaction of steps to overcome energy barrier. • There is usually what is called rate determining step. • Determining the rate of the reaction is what is called KINETICS.

  11. Metals • Corrosion is a chemical reaction. • What is considered in corrosion? • Feasibility. • Rate. • The answer is both. • You can even protect metals if you impair the feasibility or slow down the rate.

  12. Metals • Now consider the structure of metal atoms and their mutual relation. • Metal atoms have some free electrons. • In the matrix of metals, free electrons do not relate specifically to a certain atom. • They are just Free ELECTRONS. • This is what provides the metals with their Characteristic Properties.

  13. Metals • If the metals are going to react, What is better for the atoms? • To share these electrons. • To loss these electrons. • The answer, for sure is the second option. • So they are going to exchange ions with other reactants. • Ionic Bond.

  14. Corrosion Potential • Now, reaction of metals depends on DG. • But this reaction involves transfer of electrons. • DG is a measure of chemical energy. • But this energy is transferred to electric energy for the reaction to take place. • Is it possible to measure this energy using the electric parameters.

  15. Corrosion Potential • Now Recall the definition of volt. • The voltage between two points is 1 volt if the amount of energy required to transfer 1 unit charge is 1 unit of energy. • This is why P = V I • Or Energy/ time = Volt X (Charge/ Time) • Energy = Volt X Charge.

  16. Normality • If you dissolve one mole (atomic weight) in one liter of water, concentration is called morality (1 molar solution). • If you dissolve one equivalent weight in one liter of water, concentration is called Normality (1 Normal solution). • So, for one and the same material • 1 N solution = 1M solution if n =1. • 1 N solution = ½ M solution if n = 2.

  17. Normality • For explanation let us consider sodium chloride (NaCl) • At. Wt. for Na = 23 • At. Wt. for Cl = 17. • At. Wt. for NaCl = 40. • 1 M solution means 40 gram NaCl in 1 liter water. • 1 N solution is the same. • Concentration of NaCl in sea water is about 3.5% (35 grams in 1 liter).

  18. Normality • Another example is (FeCl2) • At. Wt. for Fe = 56 • At. Wt. for Cl = 17. • At. Wt. for FeCl2 = 90. • 1 M solution means 90 grams FeCl2 in 1 liter water. • 1 N solution means 45 grams FeCl2 in 1 liter water. • Normality is more expressive than molality.

  19. Cu –ve potential Reduction prevails (Cathode) Cu +ve potential Oxidationprevails (Anode) Electrolytic Cell • Now consider Electro-refining of Cu.

  20. ECu = 0.34 V Reduction prevails Positive Electrode (Cathode) EZn = -0.76 V Oxidation prevails Negative Electrode (Anode) Galvanic Cell • Now consider a cell having Cu in 1N CuSO4 solution in half cell and Zn in 1N ZnSO4 solution in the other half cell. V Electrons Current O.C.P. = 1.1 Volts

  21. Anode/ Cathode • Now if you consider Zn Zn++ + 2e → Zn • This called reduction. Or Zn → Zn++ + 2e • This is called oxidation. • As a convention we use the potential measurement for the reduction reaction. • If V increases, DG decreases, more reduction takes place, more protection.

  22. Corrosion Potential • Back to Cu + Zn • Cu will not be dissolved (protected). Cu++ + 2e → Cu • Zn will dissolve (Corroded). Zn → Zn++ + 2e • As the difference in potential increases more current takes place but not necessarily.

  23. Corrosion Rate • As we know OC potential is higher than short circuit potential. • How much is the highest current provided by any galvanic cell, let us see.

  24. Concentration Polarization • 3- Concentration Polarization: • Consider the case of placing steel part in aerated water. • Anode: Fe dissolution • Cathode: Oxygen Reduction (Oxygen). 1- Transport of oxygen to steel by diffusion. 2- Reduction of Oxygen O2 + 2H2O +4e → 4OH- Steel Low diffusion

  25. High Oxygen Diffusion Potential Low Oxygen Diffusion Fe Current Concentration Polarization Factors affecting this phenomenon: 1- Temperature. 2- Agitation. 3- Pressure 4- Flow Rate. 5- Concentration. O2

  26. Control of Rate • Decrease the metal conductance. • Lower the electrolytic conductance. • Control one of the surfaces (Larger Anode is Better. • Control one of the two reactions (anodic and Cathodic).

  27. Control of Cathodic Reactions 2H+ + 2e → H2 • Increase or Control pH. • Increase Pressure (not a solution) • Decrease Pubbling rate (not a solution in tanks and pipelines.

  28. Control of Cathodic Reactions O2 + 2 H2O + 4 e → 4 OH- • Increase or Control pH. • Use Scavengers. • In open vessels, temperature lowers the reaction rate. • In closed vessels, temperature increases rate.

  29. Uniform Attack • If we have one steel plate, corrosion will take place. • The anode and cathode will alternate from a point at the surface to another. • As the polarization of hydrogen increases at a certain point. • The other point will act as a cathode. • Uniform corrosion is not very severe usually.

  30. Measurement of Corrosion • Conversion of Current to Corrosion Rate: • If i A/cm2 is the current density, i Coulombs/cm2 transfer per second. • OR (iX365X60X60X24) = (31,536,000Xi) Coulombs per year. • OR 31,536,000 (i/96,500) = 326.79 (i) Farads per year per cm2. • If the equivalent weight of the metal is (EW), this means that (326.79XiXEW)gms/year/cm2 • If the density of the metal is (r) grams/cm3, this means that (326.79XiXEW/r) cm3 of metal corrode in one year from 1cm2 OR the metal loses (326.79XiXEW/r)cm/year. • Metal loses (13.617/2.54)X(iXEW/r) or (128.66XiXEW/r)in./year. • This means that the metal lost (128,660XiXEW/r)mils/year {mpy}. NOTE THAT THIS IS VALID ONLY FOR UNIFORM CORROSION

  31. Galvanic corrosion • If you place two dissimilar metals beside each other, the more negative potential will corrode. • Corrosion effect will increase as • Ratio of anode to cathode decreases. • Resistance of electrolyte deceases. • Criticality of corroded part increases. • Some notes about painting.

  32. Electrochemical Series • Metals are ranked in accordance with their potential in 1 N solution of their solutions. • Hydrogen is zero reference: 2H+ + 2e → H2 • If metal has positive value (Au, Ag, Pt), it is called noble metal or semi-noble (Cu). • If metal has negative value (Fe, Al, Mg, Zn), it is called active metal. • Such Ranking is called Electrochemical Series.

  33. Galvanic Series • From all the discussion, it can be noticed that every metal will have different potentials in different media. • The behavior in different media depends on many different correlated factors. • This is why Electorchemical series can be not indicative of the corrosion state. • So, Galvanic Series is more practical.

  34. Passivity • But what about if a product of corrosion is formed. • The rate of generation of product increases with current. • At a certain amount of product, it could hinder ions from dissolution into solution. • This makes the rate of corrosion very slow. • This takes place in a few metals only.

  35. Pitting and Crevice • At a certain value passivity breaks down to start the transpassivity stage. • The presence of chloride ions was found to decrease as the chloride content in the solution increases. • Chloride ions were expected to attack the passive layer leaving unprotected area. • This case represents high cathode to anode area.

  36. Pitting and Crevice • As resistance of the material increases the Epit is expected to increase. • So it can be taken as a measure of resistance to pitting. • Chromium is added to iron to increase passivity. • At 12% Cr the surface is expected to be covered with Cr2O3. • However further increase in Cr will increase the passive layer thickness and increase resistance to damage by chloride ions.

  37. Pitting and Crevice • As a rule of thumb those steels covered with 100% passive layer are called stainless steels. • Cr and Mo increases both thickness and stability of passive layer. • However, Fe++ formed in the pit will hydrolyze according to the reaction: Fe++ + H2O + 2Cl- → Fe(OH)2 + 2HCl HCl is a strong acid leading to decrease of the pH.

  38. Fe+++ Transported Cl- Fe++, H+, Cl- Probability Pitting and Crevice P/D Surface Area

  39. Pitting and Crevice • This is why pits more corrosive environment takes place within pits as they grow leading to autocatalytic action. • Nitrogen in steel was found to react with H+ in the pits to form NH3 and reduce the autocatalytic action. • For stainless steels, pitting resistance equivalent number (PREN) is equal to: PREN = Cr + 3.3 (Mo + 0.5 W) + 16N

  40. Pitting and Crevice • How to measure the resistance of material to pitting: 1- PREN will identify the grade of stainless steel. 2- Pitting potential. 3- the potential at which the anodic polarization curve intersects with the passive zone again (Eprot). • However, the difference of Epit-Eprot is more indicative of the resistance.

  41. Pitting and Crevice • Pitting is expected to grow more downward or at the upstream especially encountering elbows. • Now how to measure the intensity of pitting: • Density. • Diameter. • Depth. • Pitting factor is a measure of the prevailage of pitting against general corrosion • P/d tends to zero for general corrosion.

  42. Pitting and Crevice • P-d could be measured by: • 1- Metallography. • 2- Machining • 3- Micrometer. • 4- Microscopy. d P

  43. Pitting and Crevice

  44. Pitting and Crevice

  45. Pitting and Crevice • Crevice attack is similar to pitting in a way or another. • Inside the crevice lack of oxygen and increase in the chloride content take place leading to break down of passivity. • Thermal insulation and carbonate deposits may lead to the dame situation. • Filliform corrosion is an example of crevice attack.

  46. Pitting and Crevice Inert Washer Stainless Steel

  47. Pitting and Crevice • Even in bolts, which after rain contains some corrosive media in their crevices (does not dry easily). • Solutions include: • 1- Use larger and less number of bolts. • 2- Tighten the bolts as possible. • 3- Use ductile caulking. • 4- Use sealing compounds. • 5- Use Weathering Steel (A HSLA containing copper, phosphorous and nickel in controlled amounts).

  48. Differential Aeration • can take place even if there is no passivity. • Consider immersion of pipe in the earth (soil to air interface). • There will be difference in the mixed potential due to different values of oxygen Concentration.

  49. Pourbaix Diagram

  50. H2 Hg Platinized Platinum H2SO4 H2 Reference Electrodes • Our zero arbitrary reference electrode. • Potential =0 at STP. H+ + 2e → H2 Standard Hydrogen Electrode (SHE)