Environmental Interactions

Transcription

1 Environmental Interactions Chemical reaction between the material and its environment Beneficial interactions: materials processing Carburization and nitriding hardens for wear resistance Doping adds electrically active species Interfacial compounds are used as diffusion barriers Harmful interactions Oxidation Materials burn slowly at high T Corrosion Electrochemical reactions oxidize near room temperature

2 Oxidation Îgº (J/mole) Cu O 2 FeO NiO Cr O 2 3 SiO 2 Al O T (ºC) Almost all metals oxidize spontaneously Sparks from machining (Fe slivers burning in air) Tarnish on surfaces (may be S as well as O) Au is the great exception (hence its monetary value) Thermodynamics gives a partial explanation Free energy change for oxidation is negative (except Au) Driving force increases as T decreases But rate of oxidation is negligible except at high T

3 Kinetics of Oxidation Three Characteristic Behaviors Îm Ôt t ln(t) Parabolic growth: Δm = k p t Important at low to intermediate T Linear growth: Δm = k L t Important at high T Logarithmic growth: Δm = k ln ln(t) Observed at very low T, little engineering significance

4 Mechanism of Parabolic Oxidation O 2 O O = MO 2e - metal M ++ oxide δ δ = k t k = Aexp Q D kt Diffusion through coherent oxide film Metal is ordinarily more mobile, diffuses to oxidize at free surface Growth is diffusion controlled Thickness increases roughly as mean diffusion distance (<x> = 2Dt) Film diffusivity controls oxidation Oxidation is a high-t phenomenon (rate increases exponentially with T) Oxides with low D (high Q D ) are protective Film forms, but cannot growth

5 Protective Oxides O 2 O O = MO 2e - metal M ++ oxide δ δ = k t k = Aexp Q D kt The most stable oxides have the lowest diffusivities Metal diffusion through point defects (vacancies or interstitials) More stable oxides (low free energy) have lower defect densities The most stable oxides are formed by metals that oxidize most easily Al 2 O 3. Cr 2 O 3 are exceptionally stable Al and Cr have exceptional oxidation resistance ( fantastic finishes ) Not because they resist oxidation But because they oxidize so easily

6 Mechanisms of Linear Oxidation δ Linear oxidation is the addition of many parabolic steps Oxide does not fit perfectly on surface mechanical strain Strain increases as film thickens At critical thickness, film ruptures, exposing fresh surface Process repeats To suppress film rupture, suppress film growth Minimize diffusion through film t

7 Oxidation Resistance O 2 O O = MO 2e - metal M ++ oxide δ δ = k t k = Aexp Q D kt Thermodynamic resistance: oxide cannot form Noble metals such as Au, Pt Reliable resistance ( all the gold that ever was is somewhere today ) Kinetic resistance: oxide cannot grow Class includes the most easily oxidized metals (Al, Cr) The oxide layer forms easily, paints metal Use with caution: if oxide is penetrated, catastrophy may result

8 Breakdown of the Protective Film Al Al 2 O 3 Al-Hg (L) Al Chemical attack: Hg on Al Film broken in presence of Hg Mechanical scratch Chemical (ex: Cl from HgCl 2 ) Hg forms a liquid amalgam with Al Whiskers of oxide grow, heat Al Al catches fire Dust explosions Fine particles oxidize Oxide heats particle Film breaks down from heat Particle oxidizes rapidly (explodes) Al 2 O 3

9 Engineering Oxidation Resistance: Alloying to Create Protective Films - Stainless Steel ln (k) Cr (wt%) Influence of Cr on the oxidation rate of Fe k = Aexp Q D kt The corrosion rate of Fe decreases with Cr Asymptotes at > 8% Cr ( stainless steel ) Preferential incorporation of Cr into the oxide film Film is essentially Cr 2 O 3 when Cr >8%. Protective film no better than protective oxide Stainless steel liable to oxidation in presence of Cl (attacks Cr 2 O 3 ) Stainless steel oxidizes at sufficiently high T

10 Engineering Oxidation Resistance: Protective Coatings protective coating bonding layer protected structure Protect high temperature structures with oxidation-resistant coatings Ex: turbine blades in jet engines Properties required of a protective coating Good oxidation resistance (Al, Cr) Resistance to spall (fracture of coating) Matching coefficient of thermal expansion Intermediate bonding layer Common choices: CoCrAlY, NiCrAlY Co, Ni, Cr/Al ratio control adjust thermal expansion Y improves adhesion at interface (often add additional bonding layer

11 Fig. 1: The compressor turbine wheel and the suspect CT blade #1 Blade #1 Concave (pressure) side unstable fracture fatigue Convex (suction) side

12 Fig. 3: SEM of the fracture surface at the trailing edge of CT blade #1

13 Fig. 35: 20th section (etched) δ = 80 µm Δ = 940 µm

14 Fig. 16: 9th scetion (etched) δ = 60 µm Δ = 460 µm

15 Environmental Interactions: Aqueous Corrosion The primary source of degradation of structures Particularly steel structures ( rust ) Corrosion is a low-temperature oxidation mechanism Normal oxidation is prevented by the natural oxide coating In corrosion, the protective coating does not automatically form In the reaction: M ++ + O = = MO The metal ions form at one location (the anode ) The oxygen forms at another (the cathode ) The two do not ordinarily develop a good protective coating Corrosion is an electrochemical process Requires both electrical and chemical contact between Anode, where electrons and metal ions are generated Cathode, where electrons are consumed, O = is generated

16 Why metals don t dissolve Zn in electrolyte Dissolution creates Zn ++ ion in solution 2 electrons (e - ) in solid Free energy rises Charge imbalance Stops dissolution

17 Electrochemistry: The Galvanic cell completes the circuit Zn Zn ++ Cu + = SO 4 V = SO 4 Cu Zn Zn e - Cu + + e - Cu Zn dissolves, Cu is plated Half-cell potentials: φ Zn = φ 0 Zn + RT 2F ln[zn ++ ] φ Cu = φ 0 Cu + RT 2F ln[cu++ ] Two dissimilar metals (e.g., Zn, Cu) Connected electrically (e.g., by a wire) In contact with an electrolyte (e.g., ZnSO 4 CuSO 4 ) React according to potential (Δϕ = ϕ Cu - ϕ Zn ) Δϕ > 0 Zn + 2Cu + Zn ++ + Cu Complete circuit permits dissolution Electrons swept from anode (Zn) to cathode (Cu) Ions (SO 4= ) swept from cathode to anode

18 The Potential Difference E F eφ A eφº eφ B E F Zn e - Zn ++ electrolyte e - Zn ++ ϕº is related to the work function Potential difference between the Fermi level and free space ϕº is also affected by nature of the solution Attractive interactions with solution (Zn ++ favored) lower ϕº Repulsive interactions (Zn ++ opposed) raise ϕº Δϕº reflects the difference in Fermi level The difference in Fermi level in the free metals Interactions with the electrolyte Zn Zn e -

19 Half-Cell Reactions in Distilled water Electrode reaction Au Au e - Δϕ (25ºC) 1.50 volts Cu Cu + + e H 2 2H + + 2e Fe Fe e Cr Cr e Zn Zn e Al Al e For the reaction M 1 + M 2 + = M M 2 the potential is Δϕ 12 = Δϕ 1 Δϕ 2 The reaction proceeds if Δϕ 12 <0 Potentials in distilled water Potentials vary with electrolyte

20 The Galvanic Series in Seawater Increasingly anodic Tin Magnesium Nickel Magnesium alloys Brasses (Cu-Zn) Zinc Copper Aluminum Bronzes (Cu-Sn) Al-Cu alloys Silver solders Mild steel Nickel (passive) Wrought iron Monel (70Ni-30Cu) Cast iron Titanium 18Cr-8Ni stainless steel (non-passivated) 18Cr-8Ni stainless steel (passive) 50Pb-50Sn solder Gold Lead Increasing cathodic The more anodic material in the couple is corroded Note: alloys are (generally) cathodic to pure metals Free energy decreases on alloying

21 Concentration Cell Zn Zn e - φ Zn = φ 0 Zn + RT 2F ln[zn ++ ] Δφ = RT 2F ln Zn ++ [ ] 1 [ Zn ++ ] 2 Let a cell have Zn at both electrodes If the Zn concentration is different A potential difference is developed The side with the lower Zn concentration has lower potential Lower Zn is the anode; is corroded

22 Cathode Reactions in Fe Corrosion V Anode reactions: Fe Fe e - Fe Fe ++ Fe ++ H 2 O Fe OH - OH - Cathode reactions: Normal cathode reaction: 2e - + 1/2O 2 + H 2 O 2(OH) - Acidic solution: 2e - + 2H + H 2 Strong potential: 2e - + H 2 O 1/2H 2 + (OH) - Oxidation reaction: Fe (OH) - FeO + H 2 O Note FeO may not coat surface

23 Galvanic Couples Dissimilar metal contact Any two dissimilar conductors constitute a galvanic couple Microstructural heterogeneities More stable grain or region (lowest free energy) is the cathode High free energy due to: Mechanical deformation (defects) Chemical heterogeneity Phase or microstructural difference

24 Oxygen Concentration Cells Cathode: 2e - + 1/2O 2 + H 2 O 2OH - φ = ϕº+ RT nf ln [ O 2] 1/ 2 [ OH ] 2 Water immersion [O 2 ] decreases with depth Cathode at surface Anode at depth Corrosion below water line Pitting corrosion O 2 denuded at base of pit O 2 replenished at surface Anode at pit base Corrosion deepens pit

25 Crevice Corrosion Cathode: 2e - + 1/2O 2 + H 2 O 2OH - φ = ϕº+ RT nf ln [ O 2] 1/ 2 [ OH ] 2 Oxygen concentration cells develop at crevices Rapid crevice corrosion Attacks rivets, screw heads, etc.