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CHEM1612 - Pharmacy Week 10: Corrosion/Batteries Dr. Siegbert Schmid School of Chemistry, Rm 223 Phone: 9351 4196

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Presentation on theme: "CHEM1612 - Pharmacy Week 10: Corrosion/Batteries Dr. Siegbert Schmid School of Chemistry, Rm 223 Phone: 9351 4196"— Presentation transcript:

1 CHEM Pharmacy Week 10: Corrosion/Batteries Dr. Siegbert Schmid School of Chemistry, Rm 223 Phone:

2 Unless otherwise stated, all images in this file have been reproduced from: Blackman, Bottle, Schmid, Mocerino and Wille, Chemistry, John Wiley & Sons Australia, Ltd ISBN:

3 Lecture Electrochemistry Blackman, Bottle, Schmid, Mocerino & Wille: Chapter 12, Sections 4.8 and 4.9 Key chemical concepts: Redox and half reactions Cell potential Voltaic and electrolytic cells Concentration cells Key Calculations: Calculating cell potential Calculating amount of product for given current Using the Nernst equation for concentration cells NaCl

4 Lecture Al is an expensive metal because of the stability of its oxide Al 2 O 3. Al cannot be electrolysed from solution because H 2 O is preferentially reduced (E 0 Al = V; E H 2 0 = V). Al cannot be electrolysed from the pure oxide because it melts at too high a temperature (2045 ºC). In 1886, Hall and Herault independently developed a method for electrolytic production of Al metal, that is still used today. Hall-Herault process: dissolve Al 2 O 3 in hot cryolite, Na 3 AlF 6, which reduces the melting point to about 900 ºC. Production of Aluminium

5 Lecture At the anode graphite is oxidised to CO 2 (as a result the electrodes are rapidly used up), and fluoro-oxy ions are transformed in Al fluorides. Very high currents are used (~250,000 A) on an industrial scale. Production of Aluminium Hall-Herault process Graphite-lined furnace Figure from Silberberg, “Chemistry”, McGraw Hill, 2006.

6 Lecture Refining of Cu Electro-refining is the principle method by which Cu is refined to high purity. Less easily reduced metals remain in solution. Noble metals are not oxidised, so fall to the bottom as “mud”.

7 Lecture Corrosion: Unwanted voltaic cells The reduction of a metal oxide to a metal requires a lot of energy. This means that the reverse, oxidation of a metal to its oxide will be exothermic, and likely to be spontaneous. Metal oxide Metal Reduction + energy Oxidation, spontaneous Economically, the most important corrosion process is that of iron or steel.

8 Lecture Corrosion Corrosion is the process by which metals are oxidised in the atmosphere. In corrosion, a metal can act as both an anode and a cathode. The electrons released at the anode travel through the metal to the cathode. E o for the reaction is positive (a spontaneous process, product favoured). Al, Ti, Cr, Ni and Zn do not corrode (much) because they form an impervious oxide layer. Corrosion results in loss of structural strength. Iron roof

9 Lecture The Mechanism of Corrosion 1) Oxidation of Fe at active anode forms a pit and yields e - which travel through the metal 2) Electrons at the Fe (inactive) cathode reduce O 2 to OH - 3) Fe 2+ migrates through the drop and reacts with OH - and then O 2 to form rust.

10 Lecture Redox chemistry of corrosion The rusting of iron involves two (or more) redox reactions: Anode: 2 x {Fe  Fe e - }E ox 0 = 0.44 V Cathode: O 2 + 4H + + 4e -  2H 2 O E 0 = 1.23 V The Fe 2+ is further oxidised at the edges of the droplet, where [O 2 ] is highest: Anode: 2 x {Fe 2+  Fe 3+ + e - } E ox 0 = V Cathode: ½ x {O 2 + 4H + + 4e -  2H 2 O} E 0 = 1.23 V Iron (III) forms a very insoluble oxide (rust) which is deposited at the edge: 2Fe 3+ (aq) + (3+n) H 2 O(l)  Fe 2 O 3n H 2 O(s) + 6H + (aq)

11 Lecture You should now be able to explain some of the known features of rusting: Why does iron not rust in dry air? No water  no “salt bridge” Why does iron not rust in oxygen-free water, such as ocean depths? No oxygen  no oxidant Why does iron rust more quickly in acidic environments? H+ is a catalyst Why does iron rust more quickly at the seaside? More conductivity in the “salt bridge” Chemistry of corrosion

12 Lecture Protection against corrosion Fe can be protected by preventing O 2 and H 2 O from reaching the metal, by oiling the surface or coating with a thin film of metal oxide. Anything more readily oxidised than Fe will act as anode and prevent Fe from oxidising. These sacrificial anodes can be made of any metal that is a stronger reducing agent than Fe (“Activity Series of Metals”: Zn and Mg). This is called “cathodic protection”, and is used frequently in large iron structure such as ships, pipes, bridges, etc Zinc anode Bronze rudder

13 Lecture Mild steel bolts Stainless steel hanger Mild steel karabiner Aluminium karabiner Images from Galvanic Corrosion

14 Lecture Batteries Commercial use of redox reactions 3 classes of batteries:  Primary batteries: Non-rechargeable (e.g. alkaline battery)  Secondary batteries: Rechargeable (e.g. lead-acid, Ni-Cd, Li-ion batteries)  Fuel cells: Fuel (e.g. H 2 /O 2 ) pass through the cell, which converts chemical energy into electrical energy.

15 Lecture Primary Batteries Alkaline battery  Use a solid alkaline electrolyte paste (KOH).  Cannot be recharged, it is “dead” when its components reach equilibrium concentrations. Anode: Zn + 2OH -  ZnO + H 2 O + 2e - E ox 0 = 1.25V Cathode:2MnO 2 + H 2 O + 2e -  Mn 2 O 3 + 2OH - E 0 = 0.12V Overall:

16 Lecture Secondary Batteries Lead-acid battery (rechargeable) Used to start cars. The battery is recharged (turning it into an electrolytic cell) to re-establish non-equilibrium concentrations. Anode: Pb + HSO 4 -  PbSO 4 + H + + 2e - E ox 0 =0.30 V Cathode: PbO 2 + 3H + + HSO e -  PbSO 4 + 2H 2 O E 0 =1.63 V

17 Lecture Fuel Cells A fuel cell is a voltaic cell where the reactants are a combustible fuel, e.g. H 2, CH 4. The fuel undergoes a normal (overall) combustion reaction, however the two half-reaction are separated and the electrons harnessed. Fuel cells are still in the experimental stage, and their most notable success is probably for production of energy and water in space. Anode: H 2  2H + + 2e - ; E 0 =0.0 V Cathode: ½O 2 + 2H + + 2e -  H 2 O; E 0 =1.23 V Anode: CH 4 + 2H 2 O  CO 2 + 8H + + 8e - ;E ox 0 =-0.3 V Cathode: 4 x {½O 2 + 2H + + 2e -  H 2 O}; E 0 =1.23 V

18 Lecture Hydrogen fuel cell Pt catalyst surrounding graphite electrode ←e - Efficient No pollutants Newer designs use polymer electrolyte membrane that ferries H 3 O + groups across

19 Lecture Li-ion batteries  On discharge Li-ions move from anode to cathode.  On charge Li-ions move from cathode to anode  In the case of LiCoO 2 the battery is supplied in its discharged state.

20 Lecture Summary of Electrochemistry Concepts  Redox reactions  Standard reduction potential, E 0  Reference electrodes  Galvanic cells, cell notation, and electromotive force E cell  Electrolytic cells and Faraday’s Law  Nernst Equation and concentration cells  Examples of biological concentration cells  Relationship between E 0, ΔG, Q, and K  Corrosion  Batteries

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