Corrosion in Chemical Industry - Yildiz Technical University

YILDIZ TECHNICAL UNIVERSITY CHEMICAL ENGINEERING DEPARTMENT Corrosion in Chemical Industry SECTION 2 Prof. Dr. EMEK MÖRÖYDOR DERUN Res. Assist. ENİS MUHAMMET GÜL 1 2. ELECTROCHEMISTRY Chemical change  Electrical energy For metallic materials, the corrosion process is normally electrochemical, that...

YILDIZ TECHNICAL UNIVERSITY CHEMICAL ENGINEERING DEPARTMENT Corrosion in Chemical Industry SECTION 2 Prof. Dr. EMEK MÖRÖYDOR DERUN Res. Assist. ENİS MUHAMMET GÜL 1 2. ELECTROCHEMISTRY Chemical change  Electrical energy For metallic materials, the corrosion process is normally electrochemical, that is, a chemical reaction in which there is transfer of electrons from one chemical species to another. …………………………. 2 2. ELECTROCHEMISTRY ▪ Galvanic cell Me → Me+ + eAnode → …………… Cathode → …………… Anode (………………) A combination of two electrical conductors (electrodes) immersed in an electrolyte is called a galvanic cell. ………………………………………………………………………………………………………… 3 Electrochemical Corrosion of Metals ❖ Electrochemical corrosion is defined as the damaging of metal and metal alloys in aqueous medium. ❖ In an electrochemical system there must be ……………….transfer between anode and cathode. …………………………………………………………… ……………….. 4 Electrochemical Corrosion of Metals ▪ Zn° + 2HCl → ZnCl2 + H2  (2.1) Reaction (2.3), defined as the …………………………., is an oxidation in ▪ Zn° + 2H+ → Zn+2 + H2  (2.2) which zinc valence increases from 0 to +2, liberating electrons, e-, while (2.4), ▪ Zn° → Zn+2 + 2e- ……………………………(2.3) defined as the ……………………………, is a reduction in which the oxidation state ▪ 2H+ + 2e- → H2° …………………………………(2.4) of hydrogen decreases from +1 to 0, consuming electrons. 5 Electrochemical Cell ▪ Anode (………………………………………………………………………………) ▪ Cathode (…………………………………………………………………………) ▪ Interface (…………………………………………………………………………………………) 6 ELECTROCHEMICAL CELL ▪ Electrolytic conductor (………………………………………………………………………………………………………………) ▪ Electrolytic conductor completes the circuit. The ions in the electrolytic conductor provide the movement of current. Also it is called as ………………………… corrosion. Regional corrosion occurs in macrocell. 7 Microcell Corrosion ✓ There is potential difference between anode and cathode so H+ ions move to cathode region to be reduced. ▪ Fe → HCl Fe° + 2HCl → FeCl2 + H2 Fe° + 2H+ → Fe2+ + H2 Fe°→ Fe2+ + 2e2H++ 2e- → H2 (2.5) (2.6) (2.7) (2.8) ✓ Corrosion rate (Fe dissolution rate) increases by increasing current passing through the cell. ✓ Numerous micro corrosion regions occur on the metal surface. 8 Microcell Corrosion ▪ Corrosion shows a uniform distribution because anode and cathode replace time by time. ▪ To keep metal electrically neutral the reaction rates of equation (2.7) and (2.8) must be equal. If one of reactions is blocked the other reaction will not occur. In this situation the electronic conductive is metal. Anode reaction; ………………………….. Cathode reaction; ………………………………….(in acidic solutions) ………………………………………... Oxygen reaction in acidic solutions: ……………………………….. Neutral and alkaline solutions: ……………………………………… Metal formation: …………………………………….. 9 Metal–Ion Reduction Me+ + e- → Meo Anode: Fe → Fe+2 + 2eCathode: O2 + 4H+ + 4e- → 2H2O 10 ELECTRODE POTENTIAL Standard Electrode Electrode Reaction Potential, V 0 (V) (25°C) Hydrogen Electrode Na → Na+ + e- -2,71 Eo = 0 Mg →Mg++ + 2e- -2,37 Al →Al+3 + 3e- -1,66 Zn → Zn+2 + 2e- -0,763 Cr → Cr+3 + 3e- -0,74 Fe → Fe+2 + 2e- -0,44 Ni → Ni+2 + 2e- -0,25 Sn → Sn+2 + 2e- -0,136 Pb → Pb+2 + 2e- -0,126 H2 →2H+ + 2e- Noble metal 0 Cu → Cu+2 + 2e- 0,337 OH → ½ H2O + ¼ O2 + e- 0,40 Ag → Ag+ + e- 0,8 Pt → Pt+2 + 2e- 1,2 Au →Au+3 + 3e- 1,5 11 ELECTRODE POTENTIAL ▪ Not all metallic materials oxidize to form ions with the same degree of ease. ▪ On the left-hand side is a piece of pure iron immersed in a solution containing Fe2- ions of 1M concentration. The other side of the cell consists of a pure copper electrode in a 1M solution of Cu2+ ions. 12 ELECTRODE POTENTIAL ▪ The cell halves are separated by a membrane, which limits the mixing of the two solutions. If the iron and copper electrodes are connected electrically, reduction will occur for copper at the expense of the oxidation of iron, as follows: ▪ or Cu2+ ions will deposit (electrodeposit) as metallic copper on the copper electrode, while iron dissolves (corrodes) on the other side of the cell and goes into solution as Fe2+ ions. Thus, the two half-cell reactions are represented by the relations. 13 ELECTRODE POTENTIAL ▪ …………………………………………………………………………………………………………… …………………………………………………………………………………………………………… …………………………………………………………………………………………………………… …………………………………………………………………………………………………………… …………………………………………………………………………………………………………… ▪ An electric potential or voltage will exist between the two cell halves, and its magnitude can be determined if a voltmeter is connected in the external circuit. A potential of 0.780 V results for a copper–iron galvanic cell when the temperature is 25°C. 14 ELECTRODE POTENTIAL ▪ Various electrode pairs have different voltages; the magnitude of such a voltage may be thought of as representing the driving force for the electrochemical oxidation–reduction reaction. ▪ Metallic materials may be rated as to their tendency to experience oxidation when coupled to other metals in solutions of their respective ions. A half-cell similar to those described above [i.e., a pure metal electrode immersed in a 1M solution of its ions and at 25°C ] is termed a ………………………………… 15 3.1 Electrochemical Effects of Corrosion Thermodynamic laws G = Go + R.T.lnk G = 0 G < 0 G > 0 Go = - R.T.lnk Kinetic potential E = Eo − R.T Ink n.F The Nernst equation expresses the exact electromotive force of a cell in terms of activities of products and reactants of the cell and temperature. 16 Electrochemical Effects of Corrosion Thermodynamic laws Standard-state free energy Free energy at any moment G = Ideal gas constant Go Absolute temperature G = 0 G < 0 G > 0 Kinetic potential Cell potential + R.T.lnk Natural logarithm of the reaction coefficient Go = - R.T.lnk Standard cell potential R.T E = Eo − Ink n.F Moles of electrons Faraday's constant (96,485 coulombs per mole of electrons) The Nernst equation expresses the exact electromotive force of a cell in terms of activities of products and reactants of the cell 17and temperature. Reference Electrodes Hydrogen Electrode 18 Reference Electrodes Calomel electrode 19 THERMODYNAMICS AND POTENTIAL DIFFERENCES Corrosion Cells in Application ……………………………………………………………………………………………….. ………………………………………………………………………………………………. ……………………………………………………………………………………………….. 20 Corrosion Cells in Application ▪ When two metals electrically connected in a liquid electrolyte and create an electrochemical cell it is called as galvanic cell. The metal which has highest potential becomes an anode and corrodes, while the other acts as a cathode. Different alloys are aligned in specific environment according to their anodic or cathodic tendency. ▪ Thus, several series are obtained for sea water, clean water or industrial atmosphere. 21 CORROSION CELLS IN APPLICATION Anodic Mg Mg-alloy Zn Galvanized steel 5052 Al 3003 Al 1100 Al 6053 Al Cd 2017 Al 2024 Al Low-carbon steel Pig iron 410 stainless steel % 50 Pb-% 50 Sn (solder) Cathodic 316 stainless steel Pb Sn Cu- % 40 Zn (brass) Mn-bronze Ni based alloys AI- bronze Cu Cu-%30 Ni alloy Ni based alloys (passive) Stainless steel (passive) Ag Ti Graphite Au Pt 22 Aired salt water 23 ………………………….materials corrode rapidly than ……………………… materials. Remnants along grain boundry behaves as the cathode. (Grain boundry are more active) Pure metals have ………… corrosion resistance. 24 Roughness of Surface Coating In the process of galvanizing in which a layer of zinc is applied to the surface of steel by hot dipping, zinc is anodic to and will thus cathodically protect the steel if there is any surface damage. Any corrosion of the zinc coating will proceed at an extremely slow rate because the ratio of the anode-to-cathode surface area is quite. 25 Effects of Thermal Treatment In very low annealing temperatures corrosion rate is slow since there is only martensite phase in the steel. After tempering at the range of 200-500°C, fine precipitated ncarbide and cementite occurs. The n-carbide and cementite create the galvanic cells which accelerate the corrosion rate. Over 500°C, cementite merges with other particules and corrosion rate is reduced. 26 CONCENTRATION CELLS Salt Concentration (Metal Ions) Cell ECu2 ANODE Cu° → Cu+2 + 2e- ECu1 CATHODE Cu+2 + 2e-→ Cu° 27 Differential Aeration Cell 28 DIFFERENTIAL AERATION CELL Fe° → Fe+2 + 2e½ O2 + H2O + 2e-→ 2OH- 29 Differential Aeration Cell ❖.................................................................. Rust ❖The pipe in the sandy soil corrodes more 30 ………………………………………………………………………………………………………… Differential Temperature Cells These cells are found in heat exchangers, boilers, immersion heaters, and similar equipment. Components of these cells are electrodes of the same metal, each of which is at a different temperature, immersed in an electrolyte of the same initial composition. 31 32