1. Chapter 17 Corrosion and Degradation of Materials

2. Introduction Deterioration of a material due to chemical attack by the environment –Typically, a non-stress related failure There are exceptions Terminology –Metals: Corrosion –Dissolution in the surrounding medium –Primarily an electrochemical phenomenon Oxidation – formation of oxides –A chemical reaction –Ceramics: Chemical reaction Corrosion Oxidation of non-oxide ceramics –Polymers: Degradation –Chemical reaction with solvents –Decomposition due to UV radiation

3. Corrosion of Metals An electrochemical process, namely a chemical reaction that also involves the transport of electrons Dissolution of Zn in HCl is an electrochemical process Zn + 2HCl  ZnCl 2 + H 2 Takes place in two stages Zn  Zn 2+ + 2e - Oxidation half reaction at anode 2H + + 2e -  H 2  Reduction half reaction at cathode –Electrons that are released in the first reaction may travel through the metal and the second reaction may take place at a different location –Both oxidation and reduction reaction are necessary for corrosion to occur –Anode: location at which electrons are produced –Cathode: location at which electrons are consumed

4. Oxidation and Reduction Reactions The oxidation reaction at the anode is typically M 1  M 1 n+ + ne - The reduction reaction depends upon the environment –When H + ions are present 2H + + 2e -  H 2 –When O 2 and H + ions are present O 2 + 4H + + 4e -  2H 2 O –When O 2 and H 2 O are present O 2 + 2H 2 O + 4e -  4(OH - ) –In the presence of metal ions M 2 n+ + e -  M 2 (n-1)+

5. Electrochemical Cells An electrochemical half cell is typically a metal in contact with an electrically conductive liquid, typically a solution containing its ions Two half cells make up an electro chemical cell Depending on the metals, oxidation will occur in one cell and reduction in the other anode cathode Fe  Fe 2+ + 2e - Zn  Zn 2+ + 2e - Fe 2+ + 2e -  FeCu 2+ + 2e -  Cu The membrane keeps the solutions separated from each other, but is permeable to ions

6. Electrochemical Cells If the two electrodes (anode and cathode) are connected externally with a wire, electrons will flow from the anode to the cathode through the wire An ammeter (low resistance) will measure a current A volt meter (high resistance) will measure a potential difference between the two cells This potential difference depends upon the two half cells

7. Standard emf Series The standard Hydrogen half cell –platinum electrode –1.0M solution of H + ions –Hydrogen gas is bubbled at a pressure of 1 atmosphere This cell can be compared with other combinations of metals in contact with their ions, or in general any oxidation or reduction reaction that involves the transfer of electrons A standard half cell consists of a metal in contact with a 1.0M solution of its ions When compared with other half cells, the standard hydrogen half cell may be anodic (H 2 forms H + ions) or cathodic (H + ions combine with e - to form H 2 ) Electrons will either flow to or from the platinum electrode Other half cells will either be positive or negative with respect to the standard hydrogen half cell 1.0M  1mol of ions dissolved in 1 liter (1000 cm 3 ) of water

8. Standard emf Series The salt bridge provides a conduit for ions to flow from one solution to the other In this figure, the salt bridge is made up of KCl. Cl- and K+ ions flow through the bridge in opposite directions carrying negative and positive charge, respectively

9. Standard emf Series

10. All reactions are shown as reduction reactions To form an electrochemical cell, an oxidation and a reduction reaction are required The oxidation reaction will have a potential equal to the negative of the V 0 shown For a reduction reaction (at the cathode) M n+ + e -  M (n-1)+ V 0 For the oxidation reaction (at the anode) M (n-1)+  M n+ + e - -V 0

11. Electrochemical Cells If any two half cells are combined to form an electrochemical cell –The half cell with the more positive V 0 will be the cathode Reduction occurs at this electrode –The half cell with the more negative V 0 will be the anode Oxidation occurs at this electrode –The potential difference that is measured between the two electrodes is ∆V = V 0 cathode – V 0 anode

12. Electrochemical (Galvanic) Cells The cell shows a Zn electrode in a 1M solution of ZnSO 4 and a Cu electrode in a 1M solution of CuSO 4 During the electrochemical reaction –Zn goes into solution in the Zn half cell –Cu gets deposited on the electrode in the Cu half cell –Electrons flow from the anode (Zn) to the cathode through the wire connecting the two electrodes –Within the cell, SO 4 2- ions flow through the permeable membrane Over a period of time, the concentration of Zn in the Zn half-cell will increase while the concentration of Cu in the Cu half-cell will decrease. This will change the potential measured from the standard -1.10V

13. Cell Potentials Under standard conditions (oxidation)M 1  M 1 n+ + ne - -V 1 0 (reduction) M 2 n+ + ne -  M 2 V 2 0 Overall cell potential ∆V = V 2 0 -V 1 0 Under non-standard conditions, electrode potentials may be different –Concentration of ions different from 1M –Temperature different from room temperature Nernst Equation

14. [M 1 n+ ] and [M 2 n+ ] are the concentrations of the two ions n is the number of electrons involved in the reaction Faraday’s constant F is the charge on 1 mole of electrons F = (6.023 x 10 23 electrons/mol) x (1.602 x 10 -19 C/electron) = 96,500 C/mol R is the universal gas constant T is temperature in (K)

15. Corrosion of single electrode Single metal electrodes corrode under certain circumstances due to microscopic galvanic cells that are set up when the metal is in contact with the electrolyte Examples –Zn in contact with HCl Zn  Zn 2+ + 2e - 2H + + 2e -  H 2  –Overall reaction Zn + 2HCl  ZnCl 2 + H 2

16. Corrosion of single electrode Rusting of Fe in oxygenated water –Rust is ferric hydroxide Fe(OH) 3 –At the local anode Fe  Fe 2+ + 2e - –At the local cathode O 2 + 2H 2 O + 4e -  4(OH - ) –Overall reaction 2Fe + O 2 + 2H 2 O  2Fe(OH) 2  –In the presence of additional O 2 2Fe(OH) 2 + H 2 O + ½ O 2  2Fe(OH) 3 ( rust)

17. Examples of Galvanic cells Identify anodic and cathodic reactions and determine if corrosion will occur –Cu and Zn immersed in dilute solution of CuSO 4 Anode: Cathode: Cell potential –Cu in oxygenated water Anode: Cathode: Cell potential:

18. Examples of Galvanic cells –Mg and Fe connected by an external wire and immersed in 1% NaCl solution Anode: Cathode: Cell potential:

19. Ion Concentration Galvanic Cell Fe electrodes in contact with solutions with different concentrations of Fe 2+ ions In both cells, if the anodic reaction occurs, Fe will go into solution Fe  Fe 2+ + 2e - Use Nernst equation to calculate cell potentials 0.001M Fe 2+ 0.01M Fe 2+

20. Oxygen Concentration Galvanic Cell Differences in oxygen concentration can set up galvanic cells Fe  Fe 2+ + 2e - 0.440V O 2 + 2H 2 O + 4e -  4(OH - ) -0.401V ∆V = 0.039V Fe will go into solution in the low oxygen content solution Application –If an iron rod is left immersed in water, corrosion occurs not a the surface (high oxygen) but below the surface (low oxygen) Low O 2 conc. High O 2 conc.