CHAPTER 9: E;ECTRIC CELLS 4 U CHEMISTRY KEY IDEAS P 650 - introduction to electric cells and their uses. 9.1: Oxidation and Reduction - reduction – metals from ores see ex p 652 - reducing agents – cause ores to be reduced to metals ex: CO, C and H2 - oxidation – metals to metal compounds , first with oxygen - oxidizing agents – cause metals to be made into metal compounds, ex: O2 Cl2 Br2 p 653 practice # 1, 2, 3, 4. Electron Transfer Theory - chemical reactions involve transfer of e- , usually to make the molecules more stable - single displacement reactions easy to split into half reactions - Reduction: e- gained, oxidation # goes down Ag+1 (aq) + e- Ag (s) - Oxidation: e- lost, oxidation # goes up Cu (s) Cu+2 (aq) + 2 e- loss of e- and gain of e- happen simultaneously - reaction must be balanced for charge as well as for # atoms - Summary p 656 p 656-7 practice # 7, 8, 9, 10, 11 Oxidation States: - net electric charge an atom WOULD have if the e- were attached to the more electronegative atom ie: if all the bonds were ionic - = oxidation # - rules: see table 1 p 658 and Summary p 659 - p 659 practice # 12, 13, 14, 15, 16, 17. Oxidation Numbers and Redox Reactions - oxidation – Ox # increases - reduction – Ox # decreases - Redox reaction – both reduction and oxidation occur simultaneously - p 662 practice # 18, 19, 20, 21. - section 9.1 p 663 practice # 1, 2, 3, 4, 5, 6, 7, 8, 10. 9.2: Balancing Redox Equations - 3 reasons for having a balanced chemical equation: ________________________________________ - _________________________________________________________________________________ Oxidation # method: the entire equation is given 1. write ox # 2. write half reactions for oxidation reaction and reduction reaction 3. balance the atoms with changed ox # 4. find total # e- being lost / gained 5. multiple each half reaction by the necessary multiple so that the e- lost = e- gained 6. put coefficients from half reactions into original equation and balance the rest by inspection Ion-Electron method: may take place in acidic or basic solutions and may not have the entire equation given: 1. separate the equation into half reactions 2. for each half reaction separately: balance the atoms being oxidized / reduced 3. balance O by adding water molecules 4. balance H by adding H ions 5. balance charge by adding e6. multiply each half reaction by the necessary factor so e- on left in one half reaction = e- on right in other half reaction 7. add the half reactions and cancel any molecules and e- that are on both sides 8. FOR BASIC reactions: add OH-1 to both sides to neutralize any H+ ions 9. cancel out any extra water molecules formed p 673 practice # 6, 7, 8 section 9.2 practice p 673 # 3, 4, 6a, 7 9.3: Predicting Redox Reactions: Oxidizing and Reducing Agents – definitions The Spontaneity Rule: a reaction will be spontaneous if the oxidizing agent on the left of the table is HIGHER than the reducing agent on the right. see table appendix C11 p Predicting Redox Reactions in Solution: - copy blue box rules p 680 - Summary p 682 p 682 practice # 25, 26 section 9.3 practice # 1, 2, 3, 4, 5, 68, 9, 10, 11. 9.4: Technology of Cells and Batteries: read for interest - electric cell:________________________________________________________________________ - battery: _________________________________________________________________________ Basic Cell Design and Properties: - electrodes: ________________________________________________________________________ - electrolyte: ________________________________________________________________________ - Electric potential difference (voltage) ___________________________________________________ - Volts: ___________________________________________________________________________ - Electric current: ___________________________________________________________________ - Amperes: ________________________________________________________________________ - Coulombs: _______________________________________________________________________ - see Summary p 687 - p 687 practice # 1, 2, 3, 4 Consumer, Commercial and Industrial Cells see table p 689 for common cell types - primary cells: ______________________________________________________________________ - secondary cells: ___________________________________________________________________ ex: lead storage battery for a vehicle see p 689 Fuel Cells: Ballard Fuel Cell Aluminum Air Cell 9.5: Galvanic ( Electrochemical) Cells: - the substances that can spontaneously react in a reaction like a single displacement reaction are separated so that the e- that are transferred in the redox rxn are forced to travel through a wire, and so are available to do work. - half cell: __________________________________________________________________________ - salt bridge: ________________________________________________________________________ - porous cup: _______________________________________________________________________ A Theoretical Description of a Galvanic Cell: - copy table 1 p 696 - cathode: reduction rxn of the strongest oxidizing agent (HIGHEST on the LEFT of the table) SOAC / GERC – SEE P 697 - anode: oxidation rxn of the strongest reducing agent ( LOWEDT on the RIGHT of the table) - copy diagram of the copper with silver cell p 697 Galvanic Cells with Inert Electrodes - used when the solution, not the metal reacts, and is just a surface to deliver electrons - ex: Carbon (graphite) rod or Platinum metal rod Summary p 700 p 700 practice # 1, 2, 3, 4, 5, 6, 7b, 8. Standard Cells and Cell Potentials - standard cell: ____________________________________________________________________ - standard cell potential Eo: ___________________________________________________________ - standard reduction potential: Ero - reference half cell: the hydrogen cell against which all other potentials are measured see diagram p 702 Measuring Standard Reduction Potentials - difference between the reduction potentials of the two half cells - see ex p 702 – 704 - Eo = Eor cathode - Eor anode - if E is +, then reaction is spontaneous with the reactions as written Summary p 705 Summary p 707 p 708 practice # 10, 11, 12. 9.5 section practice p 709 # 1, 2, 3, 4, 5, 6, 7, 8. 9.6: Corrosion: - corrosion: ________________________________________________________________________ - any metal on the reduction table below the oxygen half reaction will corrode - ex: Al . It corrodes to form Al2O3 (s) but the oxide sticks well to the metal “sealing” the metal from further oxidation Rusting of Iron: - iron, water and oxygen necessary - Fe(s) Fe+2 + 2 eand ½ O2 + H2O + 2 e- 2 OH –1 - then iron II hydroxide oxidized to iron III hydroxide and iron III oxide – rust - see diagram p 711 - rust forms faster in acid rain – standard redox reaction with hydrogen ions - electrolytes speed up rusting ex: _____________________________________________________ Corrosion Prevention: ex: _______________________________________________________________________________ - coating with tin: if tin is broken, cell set up with tin and iron and the iron oxidezes, faster than without the tin - Galvanizing: coat iron with Zinc – if zinc breaks, zinc oxidizes and iron is protected Cathodic Protection: - oxidation – loss of e-, happens at the anode - to protect iron – force it to accept e- by attaching the neg end of a generator or battery to it - need constant electrical supply for it to work - used for pipelines and culverts - sacrificial anode: a metal that oxidizes more readily than iron - ex: Zn or Mg - anode metal gradually oxidizes but iron is preserved - Zn used to preserve ship hulls, Mg used to preserve hot water tanks in houses 9.6 practice p 714 # 1, 3, 4, 7, 8. Problem Set: p 726 you must do # 2a, 3, 4ace, 6, 7ac, 8ace, 10ace, 13, 14, 15ac, 18, 20b, 22, 25a. you should be able to do # 1, 2bc, 4bd, 7b, 8bd, 10bdf, 11a, 12, 15b, 16, 17, 20a, 21,