CHAPTER 23: CORROSION

Transcription

1 CHAPTER 23: CORROSION AND DEGRADATION CORROSION Corrosion: The corrosion is the deterioration of the metals (oxidation) in different environments in presence of oxygen. Anodic reaction: M M +n + ne - Cathodic reaction: 2 + 2e - H 2 (acid environment) O 2 + 2H 2 O + 4e - 4OH - (basic environment)

2 Example: CORROSION OF ZINC IN ACID Two reactions are necessary: -- oxidation reaction: Zn Zn e -- reduction reaction: 2 + 2e H 2 (gas ) Zinc flow of e - 2e in the metal - oxidation reaction Zn Zn 2+ H2(gas) Acid solution reduction reaction Other reduction reactions: -- in an acid solution -- in a neutral or base solution O e 2H 2 O O 2 + 2H 2 O + 4e 4(OH) more cathodi ic STANDARD EMF SERIES EMF series metal V o metal Au V Cu Pb Sn ore ano odic m Ni Co Cd Fe Cr Zn Al Mg Na K Metal with smaller o V metal corrodes. Ex: Cd-Ni cell o DV = V Cd 25 C Ni M 1.0 M Cd 2+ solutionni 2+ solution

3 This is how we measure the potentials in the EMF Series: Two outcomes: --Metal sample mass ne - metal, M e - M n+ ions e - H2(gas) 2e - Platinum 25 C 1M M n+ sol n 1M sol n --Metal sample mass metal, M ne - M n+ + ions H e - e - 2e - Platinum 25 C 1M M n+ sol n 1M sol n --Metal is the anode (-) --Metal is the cathode (+) o o V metal < 0 (relative to Pt) V metal > 0 (relative to Pt) Standard Electrode Potential EFFECT OF SOLUTION CONCENTRATION Ex: Cd-Ni cell with Ex: Cd-Ni cell with standard 1M solutions non-standard solutions o V Ni o VCd = o V Ni V Cd = V Ni o RT VCd nf ln X Y n = #e - per unit oxid/red Cd 25 C Ni Cd T Ni reaction (=2 here) F = 1.0 M 1.0 M X M Y M Faraday's Cd 2+ solution Ni 2+ solution Cd 2+ solution Ni 2+ solution constant =96,500 Reduce VNi -VCd by C/mol. --increasing X --decreasing Y

4 Nernst equation is used to measure the potential difference: Where: V = V 0 + RT ln nf V = New potential of half-cell, V V 0 = Standard potential, V C ion R = Real gas constant, J/mol K T = Temperature, K n = Number of electrons transferred. F = Faraday constant, 96,500 C/mol or A s/mol C ion = Molar concentration of ions If we need to measure the amount of mass lost by corrosion we use Faraday equation: Where: w = lost weight, g ItM w = = nf iatm nf M = Atomic mass of the metal, g/mol t = time, sec I = current density, A/cm 2 n = number of electrons transferred F = Faraday constant, 96,500 C/mol or A.s/mol A = area, cm 2

5 Example: Corrosion in a grapefruit Cathode Anode - Cu+ Zn H+ Zn 2+ reduction 2 + 2e H 2 (gas ) O e 2H 2 O 2e - Acid oxidation GALVANIC SERIES Ranks the reactivity of metals/alloys ll in seawater dic more catho (in nert) c m more anodic (acti ive) Platinum Gold Graphite Titanium Silver 316 Stainless Steel Nickel (passive) Copper Nickel (active) Tin Lead 316 Stainless Steel Iron/Steel Aluminum Alloys Cadmium Zinc Magnesium

6 FORMS OF CORROSION Uniform Attack Oxidation & reduction occur uniformly over surface. Crevice Between two pieces of the same metal. Rivet holes FORMS OF CORROSION (cont.) Stress corrosion Stress & corrosion work together at crack tips. Galvanic Dissimilar metals are physically joined. The more anodic one corrodes.(see Table 17.2) Zn & Mg very anodic.

7 FORMS OF CORROSION (cont.) Selective Leaching Preferred corrosion of one element/constituent (e.g., Zn from brass (Cu-Zn)). Pitting Downward propagation of small pits & holes. FORMS OF CORROSION (cont.) Erosion-corrosion Break down of passivating layer by erosion (pipe elbows). g.b. prec. attacked zones Intergranular Corrosion along grain boundaries, often where special phases exist.

8 CONTROLLING CORROSION Self-protecting metals! --Metal ions combine with O to form a thin, adhering oxide layer that t slows corrosion. Metal oxide Metal (e.g., Al, stainless steel) Reduce T (slows kinetics of oxidation and reduction) Add inhibitors --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 (e.g., paint it!). CONTROLLING CORROSION (cont.) Cathodic (or sacrificial) protection --Attach a more anodic material to the one to be protected. e.g., zinc-coated nail e.g., Mg Anode Zn 2+ e zinc zinc - Cu wire steel 2+ 2e - 2e - see Mg Mg pipe anode steel Earth

9 Rate of oxidation Rate of oxidation and the tendency of the film to protect the metal from further oxidation A0 ρm P B ratio = are related to the relative A ρ volumes of the oxide and metal. M 0 Where: P-R ratio = Pilling-Bedworth ratio A 0 = is the molecular (or formula) weight of the oxide A M = is the atomic weight of the metal ρ 0 ρ 0 = oxide density = metal density Where: ω = weight gain per unit area t = time κ L, κ p and κ e are the constant t linear, parabolic and logarithmic respectively. C and A are constant. Oxidation rate (kinetics) Linear: Parabolic ω = κ L t L ω 2 = κ p t + C Logarithmic ω = κ e log (C t + A)

10 Example: Artificial total hip replacement X-rays of a normal hip joint and fractured hip joint Schematic diagram and x-ray of an artificial total hip replacement Artificial total hip joint replacement designs Artificial total hip replacement (cont.) Three types of biomaterials are used for hip implants: Austenitic stainless steel - 316L with low sulfur content (< wt%) and extremely low carbon Cobalt-nickel-chromium-molybdenum MP35N with a composition 35 wt% Co, 35 wt% Ni, 20 wt% Cr, and 10 wt% Mo. Titanium Ti-6Al-4V with a composition 90 wt% Ti, 6 wt% Al and 4 wt% V

11 Artificial Total Hip replacement SUMMARY Corrosion occurs due to: --the natural tendency of metals to give up electrons. --electrons are given up by an oxidation reaction. --these electrons then are part of a reduction reaction. Metals with a more negative Standard Electrode Potential are more likely to corrode relative to other metals. The Galvanic Series ranks the reactivity of metals in seawater. t Increasing T speeds up oxidation/reduction reactions. Corrosion may be controlled by: -- using metals which form -- adding inhibitors a protective oxide layer -- reducing T g p -- painting --using cathodic protection.