1 / 20

CHEM1612 - Pharmacy Week 10: Corrosion/Batteries

CHEM1612 - Pharmacy Week 10: Corrosion/Batteries. Dr. Siegbert Schmid School of Chemistry, Rm 223 Phone: 9351 4196 E-mail: siegbert.schmid@sydney.edu.au. Unless otherwise stated, all images in this file have been reproduced from:

komala
Download Presentation

CHEM1612 - Pharmacy Week 10: Corrosion/Batteries

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. CHEM1612 - PharmacyWeek 10: Corrosion/Batteries Dr. SiegbertSchmid School of Chemistry, Rm 223 Phone: 9351 4196 E-mail: siegbert.schmid@sydney.edu.au

  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. 2008      ISBN: 9 78047081 0866

  3. 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. Production of Aluminium • Al is an expensive metal because of the stability of its oxide Al2O3. • Al cannot be electrolysed from solution because H2O is preferentially reduced (E0Al = -1.66 V; EH20= -0.42 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 Al2O3 in hot cryolite, Na3AlF6, which reduces the melting point to about 900 ºC.

  5. Production of Aluminium Hall-Herault process Graphite-lined furnace Figure from Silberberg, “Chemistry”, McGraw Hill, 2006. At the anode graphite is oxidised to CO2 (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.

  6. http://electrochem.cwru.edu/ed/encycl/fig/m02/m02-f06b.jpg 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. Reduction + energy Metal Metal oxide Oxidation, spontaneous 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. • Economically, the most important corrosion process is that of iron or steel.

  8. Corrosion Iron roof • 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. • Eo 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.

  9. 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 O2 to OH- 3) Fe2+ migrates through the drop and reacts with OH- and then O2 to form rust.

  10. The Fe2+ is further oxidised at the edges of the droplet, where [O2] is highest: Anode: 2 x {Fe  Fe2+ + 2e-}Eox0 = 0.44 V Cathode: O2+ 4H+ + 4e-  2H2O E0 = 1.23 V Anode: 2 x {Fe2+ Fe3+ + e-} Eox0 = -0.77 V Cathode: ½ x {O2 + 4H+ + 4e-  2H2O} E0 = 1.23 V Iron (III) forms a very insoluble oxide (rust) which is deposited at the edge: 2Fe3+(aq) + (3+n) H2O(l)  Fe2O3•n H2O(s) + 6H+(aq) Redox chemistry of corrosion • The rusting of iron involves two (or more) redox reactions:

  11. Chemistry of corrosion 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”

  12. Protection against corrosion • Fe can be protected by preventing O2 and H2O 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. Galvanic Corrosion Stainless steel hanger Aluminium karabiner Mild steel bolts Mild steel karabiner Images from http://www.theleedswall.co.uk/ymc/boltfund.htm

  14. 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. H2/O2) pass through the cell, which converts chemical energy into electrical energy.

  15. Anode: Zn + 2OH- ZnO + H2O + 2e-Eox0= 1.25V Cathode: 2MnO2 + H2O + 2e-  Mn2O3 + 2OH-E0= 0.12V Overall: Primary Batteries • Alkaline battery • Use a solid alkaline electrolyte paste (KOH) . • Cannot be recharged, it is “dead” when its components reach equilibrium concentrations.

  16. Anode: Pb+ HSO4- PbSO4 + H+ + 2e-Eox0=0.30 V Cathode: PbO2 + 3H+ + HSO4- + 2e-  PbSO4 + 2H2O E0=1.63 V 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.

  17. Anode: H2 2H+ + 2e-; E0=0.0 V Cathode: ½O2 + 2H+ + 2e-  H2O; E0=1.23 V Anode: CH4 + 2H2O  CO2 + 8H+ + 8e-; Eox0=-0.3 V Cathode: 4 x {½O2 + 2H+ + 2e-  H2O}; E0=1.23 V Fuel Cells • A fuel cell is a voltaic cell where the reactants are a combustible fuel, e.g. H2, CH4. 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 inspace .

  18. Hydrogen fuel cell • Efficient • No pollutants • Newer designs use polymer electrolyte membrane that ferries H3O+groups across ←e- Pt catalyst surrounding graphite electrode

  19. 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 LiCoO2 the battery is supplied in its discharged state.

  20. Summary of Electrochemistry Concepts • Redox reactions • Standard reduction potential, E0 • Reference electrodes • Galvanic cells, cell notation, and electromotive force Ecell • Electrolytic cells and Faraday’s Law • Nernst Equation and concentration cells • Examples of biological concentration cells • Relationship between E0, ΔG, Q, and K • Corrosion • Batteries

More Related