1. Chemical vs. Electrochemical Reactions  Chemical reactions are those in which elements are added or removed from a chemical species.  Electrochemical reactions are chemical reactions in which not only may elements may be added or removed from a chemical species but at least one of the species undergoes a change in the number of valance electronS.  Corrosion processes are electrochemical in nature.

2.  Chemical corrosion: Removal of atoms from a material by virtue of the solubility or chemical reaction between the material and the surrounding liquid. EXAMPLES:  Dezincification: A special chemical corrosion process by which both zinc and copper atoms are removed from brass, but the copper is replated back onto the metal.  Graphitic corrosion: A special chemical corrosion process by which iron is leached from cast iron, leaving behind a weak, spongy mass of graphite. 2 Chemical Corrosion

3. 3 ©2003 Brooks/Cole, a division of Thomson Learning, Inc. Thomson Learning ™ is a trademark used herein under license. Photomicrograph of a copper deposit in brass, showing the effect of dezincification (x50).

4. 4  Electrochemical corrosion - Corrosion produced by the development of a current in an electrochemical cell that removes ions from the material.  Electrochemical cell - A cell in which electrons and ions can flow by separate paths between two materials, producing a current which, in turn, leads to corrosion or plating.  Oxidation reaction - The anode reaction by which electrons are given up to the electrochemical cell.  Reduction reaction - The cathode reaction by which electrons are accepted from the electrochemical cell. Electrochemical Corrosion

5. 5 ©2003 Brooks/Cole, a division of Thomson Learning, Inc. Thomson Learning ™ is a trademark used herein under license. The components in an electrochemical cell: (a) a simple electrochemical cell and (b) a corrosion cell between a steel water pipe and a copper fitting.

6. 6 The anode and cathode reactions in typical electrolytic corrosion cells: (a) the hydrogen electrode, (b) the oxygen electrode, and (c) the water electrode.

7. Electrochemistry Thermodynamics at the electrode

8. Redox (Review)  Oxidation is...  Loss of electrons  Reduction is...  Gain of electrons  Oxidizing agents oxidize and are reduced  Reducing agents reduce and are oxidized

9. Redox Review (Cu-zn)  Zn displaces Cu from CuSO 4 (aq)  In direct contact the enthalpy of reaction is dispersed as heat, and no useful work is done  Redox process:  Zn is the reducing agent  Cu 2+ is the oxidizing agent

10. Separating the combatants  Each metal in touch with a solution of its own ions  External circuit carries electrons transferred during the redox process  A “salt bridge” containing neutral ions completes the internal circuit.  The energy released by the reaction in the cell can perform useful work – like lighting a bulb

11. Labelling the parts

12. Cell notation  Anode on left, cathode on right  Electrons flow from left to right  Oxidation on left, reduction on right  Single vertical = electrode/electrolyte boundary  Double vertical = salt bridge Anode: Zn →Zn 2+ + 2e Cathode: Cu 2+ + 2e →Cu

13. Odes to a galvanic cell  Cathode  Where reduction occurs  Where electrons are consumed  Where positive ions migrate to  Has positive sign  Anode  Where oxidation occurs  Where electrons are generated  Where negative ions migrate to  Has negative sign

14. The role of inert electrodes  Not all cells start with elements as the redox agents  Consider the cell  Fe can be the anode but Fe 3+ cannot be the cathode.  Use the Fe 3+ ions in solution as the “cathode” with an inert metal such as Pt

15. Anode Catho de Oxidati on Reduct ion

16. Connections: cell potential and free energy  The cell in open circuit generates an electromotive force (emf) or potential or voltage. This is the potential to perform work  Energy is charge moving under applied voltage

17. Relating free energy and cell potential The Faraday : F = 96 485 C/mole Standard conditions (1 M, 1 atm, 25°C)

18. Standard Reduction Potentials  The total cell potential is the sum of the potentials for the two half reactions at each electrode  E cell = E cath + E an  From the cell voltage we cannot determine the values of either – we must know one to get the other  Enter the standard hydrogen electrode (SHE)  All potentials are referenced to the SHE (E H =0 V)

19. Unpacking the SHE  The SHE consists of a Pt electrode in contact with H 2 (g) at 1 atm in a solution of 1 M H + (aq).  The voltage of this half-cell is defined to be 0 V.  An experimental cell containing the SHE half-cell with other half-cell gives voltages which are the standard potentials for those half-cells E cell = 0 + E half-cell

20. Zinc half-cell with SHE  Cell measures 0.76 V  Standard potential for Zn(s) = Zn 2+ (aq) + 2e : 0.76 V

21. Where there is no SHE  In this cell there is no SHE and the measured voltage is 1.10 V

22. Standard reduction potentials  Any half reaction can be written in two ways:  Oxidation: M = M + + e (+V)  Reduction: M + + e = M (-V)  Listed potentials are standard reduction potentials

24. Applying standard reduction potentials  Consider the reaction  What is the cell potential?  The half reactions are:  E° = 0.80 V – (-0.76 V) = 1.56 V  NOTE: Although there are 2 moles of Ag reduced for each mole of Zn oxidized, we do not multiply the potential by 2.

25. Extensive VS intensive  Free energy is extensive property so need to multiply by no of moles involved  But to convert to E we need to divide by no of electrons involved  E is an intensive property

26. The Nernst equation  Working in nonstandard conditions

27. 27  Electrode potential - Related to the tendency of a material to corrode. The potential is the voltage produced between the material and a standard electrode.  emf series - The arrangement of elements according to their electrode potential, or their tendency to corrode.  Nernst equation - The relationship that describes the effect of electrolyte concentration on the electrode potential in an electrochemical cell.  Faraday’s equation - The relationship that describes the rate at which corrosion or plating occurs in an electrochemical cell. Summary

28. 28 ©2003 Brooks/Cole, a division of Thomson Learning, Inc. Thomson Learning ™ is a trademark used herein under license. The half-cell used to measured the electrode potential of copper under standard conditions. The electrode potential of copper is the potential difference between it and the standard hydrogen electrode in an open circuit. Since E 0 is great than zero, copper is cathodic compared with the hydrogen electrode.

29. 29

30. 30 Suppose 1 g of copper as Cu 2+ is dissolved in 1000 g of water to produce an electrolyte. Calculate the electrode potential of the copper half-cell in this electrolyte. The atomic mass of copper is 63.54 g/mol. Example 22.1 SOLUTION From chemistry, we know that a standard 1-M solution of Cu 2+ is obtained when we add 1 mol of Cu 2+ (an amount equal to the atomic mass of copper) to 1000 g of water. The atomic mass of copper is 63.54 g/mol. The concentration of the solution when only 1 g of copper is added must be: Example HalfCell Potential for Copper From the Nernst equation, with n = 2 and E 0 = +0.34 V:

31. 31  Polarization - Changing the voltage between the anode and cathode to reduce the rate of corrosion. –Activation polarization is related to the energy required to cause the anode or cathode reaction –Concentration polarization is related to changes in the composition of the electrolyte –Resistance polarization is related to the electrical resistivity of the electrolyte. Polarization

32. 32 ©2003 Brooks/Cole, a division of Thomson Learning, Inc. Thomson Learning ™ is a trademark used herein under license. Photomicrograph of intergranular corrosion in a zinc die casting. Segregation of impurities to the grain boundaries produces microgalvanic corrosion cells (x50).

33. Corrosion Cells  Galvanic cell (Dissimilar electrode cell) – dissimilar metals  Salt concentration cell – difference in composition of aqueous environment  Differential aeration cell – difference in oxygen concentration  Differential temperature cell – difference in temperature distribution over the body of the metallic material

34. Dissimilar Electrode Cell  When a cell is produced due to two dissimilar metals it is called dissimilar electrode cell  Dry cell  Local action cell  A brass fitting connected to a steel pipe  A bronze propeller in contact with the steel hull of a ship Zn anode HCl Solution Cu cathode

35. Differential Temperature Cell  This is the type of cell when two identical electrodes are immersed in same electrolyte, but the electrodes are immersed into solution of two different temperatures  This type of cell formation takes place in the heat exchanger equipment where temperature difference exists at the same metal component exposed to same environment  For example for CuSO 4 electrolyte & Cu electrode the electrode in contact with hot solution acts as cathode.

36. Salt Concentration Cell

37. Differential Aeration Cell

38. Corrosion at the bottom of the electrical poles

39. Local Action Cell