Redox Reactions

1. Redox Reactions Oxidation-Reduction

2. Oxidation-Reduction • 4 Fe + 3 O2 2 Fe2O3 • 2 Fe2O3 + 3 C  3 CO2 + 4 Fe • Oxidation-"addition" of oxygen. • Reduction-"removal" of oxygen. • Change in oxidation state due to "loss" (+ oxidation state) or "gain" (- oxidation state) of electrons. • "Leo goes Ger" • Loss of electrons – oxidation • Gain of electrons – reduction

3. Oxidation-loss of electrons or gain of oxygen. Reduction-complete or partial gain of electrons. Reducing agent-substance that is reduces another substance by losing electrons (is actually oxidized). Oxidizing agent-substance that has oxidized another substance by accepting its electrons (is actually reduced).

4. Will have at least one type of atom releasing electrons and another type of atom accepting electrons. Label every atom on both sides of equation with its oxidation #. Look for change in oxidation # in same atom on both sides equation. Must have at least 1 atom losing/at least 1 atom gaining electrons. Metals always lose electrons. Nonmetals always gain electrons. Redox reactions

8. Volume 2 • Activity 8: Color Reactions that Involve the Transfer of Electrons (pp 566-586)

9. Homework: Read 11.1-11.2, pp. 337-342 Q pg. 358, #1-5 Q pg. 359, #3, 4, 6

10. Electrochemistry: study of how electricity produces chemical reactions and chemical reactions produces electricity. Electrochemical Cells (any device which converts chemical energy into electrical energy or vice versa)

11. Piece of copper placed into AgNO3 solution Spontaneous reaction occurs. Grayish white silver deposit formed on copper. Solution turns blue because of copper (II) nitrate: 2Ag+ + Cu  Cu2+ + 2Ag No usable energy harnessed because it is dissipated as heat. Same reaction can occur and produce electricity if placed in galvanic cell.

12. Voltaic or Galvanic Cell • Type of electrochemical cell which converts chemical energy into electrical energy. • Common batteries are one or more voltaic cells • Electrolytic cells-externally supplied electric current to drive chemical reaction which would not occur spontaneously. • External electric current flow is created using two different metals. • Metals differ in their tendency to lose electrons.

13. Voltaic cell consists of: At least two half-cells (compartments): Reduction cell Oxidation cell Salt bridge between the two halves.

14. Half cells • Each half cell consists of • Electrode. • Anode-where oxidation takes place. • Cathode-where reduction takes place. • Object that conducts electrons to or from another substance. • Electrolyte solution. • Separate oxidation and reduction reactions take place here.

15. Solution • Contains ions from electrode by oxidation and reduction reaction. • 2 containers each store half-reactions of reaction. For example, • Reduction: Ag+ + e- Ag • Oxidation: Cu  Cu2+ + 2e-

16. Wire • Carries electrons in external circuit from anode  cathode: • Electric charge can flow between two points only when a difference in electrical potential energy exists between two points, which are electrodes. • Electron flow from anode (-) to cathode (+) that creates electricity. • For voltaic cell to continue to produce external electric current-sulfate ions in solution move from right  left to balance electron flow in external circuit • Completes circuit. • Metal ions can’t move between electrodes: • Porous membrane or salt bridge provides selective movement of - ions in electrolyte.

17. Salt Bridge • Tube contains conducting solution of soluble salt (ionic compound), such as KCl, NaCl, KNO3, held in place by agar gel or other form of plug. • Completes electrical circuit but does not allow electron flow or mixing of two reacting solutions. • Ions can enter and leave solutions and move through plug but solution cannot. • Ion flow between solutions and prevent build up of charge on one side which would stop electron flow.

18. Zinc/copper metal rods (electrodes) placed in solutions of their salts. Conductor in a circuit that carries electrons to or from a substance other than a metal. Electrons flow through external wire from zinc  copper. Type reaction determines electrode type: Anode: Zinc more readily loses electrons than copper. Becomes positive ion. Goes into aqueous solution Decreases mass of Zn electrode. Cathode: Cu receives 2 electrons. Cu ion from solution  uncharged Cu atom which deposits on Cu electrode. Increases mass. Electrochemical cell created by placing metallic electrodes into electrolyte where a chemical reaction either uses/generates electric current

19. Two reactions are typically written : Zn(s)  Zn2+(aq) + 2e- Zinc "half-reaction" is oxidation (loses electrons) and occurs at anode (- terminal of battery). Cu2+(aq) + 2e-  Cu(s) Copper "half-reaction" is reduction (gains electrons) and occurs at "cathode" (+ terminal of battery). Cu2+(aq) + 2e-  Cu(s) Zn(s)  Zn2+(aq) + 2e- Zn (s) + Cu2+ (aq)  Zn+2(aq) + Cu(s)

20. Oxidation half-cell (anode) on left. Double vertical line-salt bridge/porous disc separ-ates anode compartment from cathode compartment. Single vertical line indicates boundariesof phases in contact. Zn rod/zinc sulfate solution separate phases in physical contact. Zn (s) l ZnSO4(aq) ll CuSO4(aq) l Cu(s)

22. Electricalconductor • K+ • Cl- SO4 Reducing agent Oxidizing agent Oxidation half-reaction Zn  Zn2+ + 2e- Reduction half-reaction Cu2+ + 2d-  Cu

24. Example • Write anode/cathode reactions for galvanic cell that uses reaction Ni(s) + 2 Fe3+ Ni2+ + 2 Fe2+ • Oxidation (at anode) and electrode must be Ni | Ni2+, • Ni(s)  Ni2+(aq) + 2 e- • Reduction (at cathode): Fe3+ l Fe2+: • 2 Fe3+ + 1 e-  2 Fe2+ • For every Ni atom oxidized, two Fe3+ ions reduced. • Electrons from Ni metal flow from anode, pass the load, and then carry out reduction at surface of cathode to reduce ferric ions to ferrous ions. • Ions in solution move accordingly to keep charges balanced. • Galvanic cell is: Ni(s) | Ni2+(aq) || Fe3+(aq), Fe2+(aq) | Pt(s) • Inert Platinum electrode placed in solution-provides e’s for reduction.

26. Comparison: AnodeCathode • Oxidation occursReduction occurs • Electrons producedElectrons consumed • Anions migrate towardCations migrate toward • Has negative signHas positive sign

27. Batteries

28. Anode: (negative battery pole) Zinc can/container in contact w/MnO2 (ion transfer occurs here). Electrolyte (NH3Cl-source of H+ and ZnCl2 in water). Starch added to thicken it to paste-like consistency to prevent leaks. Cathode: (positive battery pole) Carbon rod (inactive cathode-made of material that does not participate in redox reaction). Voltage produced ~1.5 V Heavy, don’t last long, cheap. Dry Cell (Leclanche cell) are voltaic cells

30. Package of 1/more galvanic cells (connected in series) used for production and storage of electric energy by chemical means. Primary battery- Produces electric energy by means of redox reactions that are not easily reversed. Deliver current until reactants are gone and then battery is discarded. Secondary battery (storage battery)- Rechargeable. Depends on reversible redox reactions Batteries

33. Anode: Zinc amalgamated w/mercury. Strong alkaline electrolyte containing Zn(OH)2 /HgO No change in electrolyte composition during operation: Overall cell reaction involves only solid substance. Provides more constant voltage (1.35 V) than dry cell. Higher capacity, longer life but more expensive than dry cell. Mercury Battery(medicine/electronics-pacemakers, hearing aids, electric watches, light meters)

34. When recharged (reverse normal electrochemical reaction by applying external voltage at cathode/anode replenishes original materials). Nonspontaneous reverse reaction driven by electricity from alternator. Could go on forever but: Lead sulfate formed in reduction falls to bottom. Lead electrodes get thinner & more brittle. Eventually break, short out cell, battery needs replacing. Lead Storage Battery

35. “Gone dead” in cold climates Electrolyte must be fully conducting to function properly. If cold, ions move more slowly through viscous medium, decreasing power output of battery. If warmed to room temperature, recovers normal power.

36. Employs solid (not aqueous solution or water-based paste) as electrolyte connecting electrodes. Ions (not electrons) migrate through solid polymer electrolyte from anode to cathode while electrons flow externally from anode to cathode to complete circuit(TiS2 or V6O13). Solid-state Lithium Battery

37. Advantage of lithium: As anode-most negative E0 value. Lightweight metal-less mass. Voltage as high as 3 V-higher capacity. Can be recharged in same way as lead storage battery. Last longer than other kinds of batteries.

39. Fuel Cells • Voltaic cells in which oxidized substance is continually supplied as fuel. • Do not store chemical energy as batteries do: • Reactants must be constantly resupplied and products must be constantly removed from fuel cell-resembles engine more than battery. • Can be 70% efficient. • Generally free of noise, vibration, heat transfer, thermal pollution and other problems associated with conventional power plants.

40. Advantages of fuel cells: No recharging. No harmful pollutants. Run more quietly/ cheaply than conven-tional internal combus-tion engines or elec-trical generators. Disadvantages: High initial cost. Fuels not readily available for conventional cars (gasoline or diesel). KOH (aq)

41. Homework: Read 11.3, pp. 342-350 Q pg. 358, #10-12 Q pp. 359-360, #7-9, 11, 13, 16

42. Corrosion… …and preventing it

44. Corrosion • Corrosion-loss of metal resulting from oxidation-reduction reaction of the metal with substances in environment. • Rust-forms on iron in presence of water and oxygen: • Fe2O3-iron(III) oxide. • Occurs quickly in salt water-electrical circuit completed by migration of electrons and ions. • Rust on surface of iron too porous to protect underlying metal.

45. Aluminum-more negative standard reduction potential than iron but corrosion doesn’t occur to the extent as iron because a layer of insoluble aluminum oxide (Al2O3) forms on surface when metal exposed to air and that protects aluminum underneath. Coinage metals also corrode, but much more slowly: Copper forms layer of copper carbonate (CuCO3)-green patina. Silver forms silver sulfide (Ag2S)-tarnish.

46. Protect metals from corrosionCoat metal surface with paint-good only as long as paint is good • Alloys of iron with other metals: • One or more elements w/characteristics of metal. • Pure metals and alloys have different physical properties. • Can be harder than pure metal (copper + gold  jewelry). • Apply thin layer of tin over iron-good as long as tin remains intact. • Galvanizing-iron is coated w/layer of zinc and even though zinc is more readily oxidized than iron, it is a self-protecting metal. • When exposed to air, zinc oxidizes at surface, thin metal oxide coating clings tightly to metal and seals it from further oxidation.

47. Anode: Water oxidized to form oxygen: 2H2O(l)  O2(g) + 4H+(aq) + 4e- Cathode: Water reduced to form hydrogen: 2H2O(l)  H2(g) + 2OH-(aq) Overall reaction:  2H2O(l)  2H2(g) + O2(g) Electrolysis of Water

48. Electroplating Using an electrolytic cell to deposit a thin layer of metal on object. Object made to be the cathode and the aqueous metal is reduced to form the elemental metal as covering on object. Process accomplished in electrolytic cells:

49. Homework: Read 351-353 (last part of 11.3) and 11.4, pp. 353-357 Q pg. 358, #13, 14, 16-18 Use link for quiz and submit by email. http://wps.aw.com/bc_suchocki_chemistry_se/0,6823,526419-nav_and_content,00.utf8.html