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Electrochemistry

AP Chemistry Chapter 20 Notes. Electrochemistry. Applications of Redox. Review. Oxidation reduction reactions involve a transfer of electrons. OIL- RIG Oxidation Involves Loss Reduction Involves Gain LEO-GER Lose Electrons Oxidation Gain Electrons Reduction. Applications.

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Electrochemistry

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  1. AP Chemistry Chapter 20 Notes Electrochemistry Applications of Redox

  2. Review • Oxidation reduction reactions involve a transfer of electrons. • OIL- RIG • Oxidation Involves Loss • Reduction Involves Gain • LEO-GER • Lose Electrons Oxidation • Gain Electrons Reduction

  3. Applications • Moving electrons is electric current. • 8H++MnO4-+ 5Fe+2 +5e-® Mn+2 + 5Fe+3 +4H2O • Helps to break the reactions into half rxns. • 8H++MnO4-+5e-® Mn+2 +4H2O • 5Fe+2® 5Fe+3 + 5e- ) • In the same mixture it happens without doing useful work, but if separate

  4. Connected this way the reaction starts • Stops immediately because charge builds up. H+ MnO4- Fe+2

  5. Galvanic Cell Salt Bridge allows current to flow H+ MnO4- Fe+2

  6. e- • Electricity travels in a complete circuit • Instead of a salt bridge H+ MnO4- Fe+2

  7. Porous Disk H+ MnO4- Fe+2

  8. e- e- e- e- Anode Cathode e- e- Reducing Agent Oxidizing Agent

  9. Cell Potential • Oxidizing agent pushes the electron. • Reducing agent pulls the electron. • The push or pull (“driving force”) is called the cell potential Ecell • Also called the electromotive force (emf) • Unit is the volt(V) • = 1 joule of work/coulomb of charge • Measured with a voltmeter

  10. 0.76 H2 in Cathode Anode H+ Cl- Zn+2 SO4-2 1 M ZnSO4 1 M HCl

  11. Standard Hydrogen Electrode • This is the reference all other oxidations are compared to • Eº = 0 • º indicates standard states of 25ºC, 1 atm, 1 M solutions. H2 in H+ Cl- 1 M HCl

  12. Cell Potential • Zn(s) + Cu+2 (aq)® Zn+2(aq) + Cu(s) • The total cell potential is the sum of the potential at each electrode. • Eºcell = EºZn® Zn+2 + EºCu+2® Cu • We can look up reduction potentials in a table. • One of the reactions must be reversed, so change it sign.

  13. Cell Potential • Determine the cell potential for a galvanic cell based on the redox reaction. • Cu(s) + Fe+3(aq)® Cu+2(aq) + Fe+2(aq) • Fe+3(aq)+ e-® Fe+2(aq) Eº = 0.77 V • Cu+2(aq)+2e-® Cu(s) Eº = 0.34 V • Cu(s) ® Cu+2(aq)+2e-Eº = -0.34 V • 2Fe+3(aq)+ 2e-® 2Fe+2(aq) Eº = 0.77 V

  14. Line Notation • solid½Aqueous½½Aqueous½solid • Anode on the left½½Cathode on the right • Single line different phases. • Double line porous disk or salt bridge. • If all the substances on one side are aqueous, a platinum electrode is indicated. • For the last reaction • Cu(s)½Cu+2(aq)½½Fe+2(aq),Fe+3(aq)½Pt(s)

  15. Galvanic Cell • The reaction always runs spontaneously in the direction that produced a positive cell potential. • Four things for a complete description. • Cell Potential • Direction of flow • Designation of anode and cathode • Nature of all the components- electrodes and ions

  16. Practice • Completely describe the galvanic cell based on the following half-reactions under standard conditions. • MnO4- + 8 H+ +5e-® Mn+2 + 4H2O Eº=1.51 • Fe+3 +3e-® Fe(s) Eº=0.036V

  17. Potential, Work and DG • emf = potential (V) = work (J) / Charge(C) • E = work done by system / charge • E = -w/q • Charge is measured in coulombs. • -w = qE • Faraday = 96,485 C/mol e- • q = nF = moles of e- x charge/mole e- • w = -qE = -nFE = DG

  18. Potential, Work and DG • DGº = -nFE º • if E º < 0, then DGº > 0 spontaneous • if E º > 0, then DGº < 0 nonspontaneous • In fact, reverse is spontaneous. • Calculate DGº for the following reaction: • Cu+2(aq)+ Fe(s) ® Cu(s)+ Fe+2(aq) • Fe+2(aq)+ e-® Fe(s) Eº = 0.44 V • Cu+2(aq)+2e-® Cu(s) Eº = 0.34 V

  19. Cell Potential and Concentration • Qualitatively - Can predict direction of change in E from LeChâtelier. • 2Al(s) + 3Mn+2(aq) ® 2Al+3(aq) + 3Mn(s) • Predict if Ecell will be greater or less than Eºcell if [Al+3] = 1.5 M and [Mn+2] = 1.0 M • if [Al+3] = 1.0 M and [Mn+2] = 1.5M • if [Al+3] = 1.5 M and [Mn+2] = 1.5 M

  20. The Nernst Equation • DG = DGº +RTln(Q) • -nFE = -nFEº + RTln(Q) • E = Eº - RTln(Q) nF • 2Al(s) + 3Mn+2(aq) ® 2Al+3(aq) + 3Mn(s) Eº = 0.48 V • Always have to figure out n by balancing. • If concentration can gives voltage, then from voltage we can tell concentration.

  21. The Nernst Equation • As reactions proceed concentrations of products increase and reactants decrease. • Reach equilibrium where Q = K and Ecell = 0 • 0 = Eº - RTln(K) nF • Eº = RTln(K) nF • nFEº = ln(K) RT

  22. Batteries are Galvanic Cells • Car batteries are lead storage batteries. • Pb +PbO2 +H2SO4®PbSO4(s) +H2O • Dry Cell Zn + NH4+ +MnO2 ® Zn+2 + NH3 + H2O • Alkaline Zn +MnO2 ® ZnO+ Mn2O3 (in base) • NiCad • NiO2 + Cd + 2H2O ® Cd(OH)2 +Ni(OH)2

  23. Corrosion • Rusting - spontaneous oxidation. • Most structural metals have reduction potentials that are less positive than O2 . • Fe ® Fe+2+2e-Eº= 0.44 V • O2 + 2H2O + 4e- ® 4OH- Eº= 0.40 V • Fe+2 + O2 + H2O ® Fe2 O3 + H+ • Reaction happens in two places.

  24. Salt speeds up process by increasing conductivity Water Rust e- Iron Dissolves- Fe ® Fe+2

  25. Preventing Corrosion • Coating to keep out air and water. • Galvanizing - Putting on a zinc coat • Has a lower reduction potential, so it is more. easily oxidized. • Alloying with metals that form oxide coats. • Cathodic Protection - Attaching large pieces of an active metal like magnesium that get oxidized instead.

  26. Electrolysis • Running a galvanic cell backwards. • Put a voltage bigger than the potential and reverse the direction of the redox reaction. • Used for electroplating.

  27. 1.10 e- e- Zn Cu 1.0 M Cu+2 1.0 M Zn+2 Cathode Anode

  28. A battery >1.10V e- e- Zn Cu 1.0 M Cu+2 1.0 M Zn+2 Cathode Anode

  29. Calculating plating • Have to count charge. • Measure current I (in amperes) • 1 amp = 1 coulomb of charge per second • q = I x t • q/nF = moles of metal • Mass of plated metal • How long must 5.00 amp current be applied to produce 15.5 g of Ag from Ag+

  30. Other uses • Electroysis of water. • Seperating mixtures of ions. • More positive reduction potential means the reaction proceeds forward. • We want the reverse. • Most negative reduction potential is easiest to plate out of solution.

  31. Balancing Redox Equations

  32. 2. in base Am3+(aq) + S2O82-(aq) ----> AmO2+(aq) + SO42-(aq)

  33. 3. MnO4-(aq) + H2C2O4(aq)  Mn2+(aq) + CO2(g)

  34. 4. Bi(OH)3 + SnO22- Bi(s) + SnO32-

  35. ELECTROLYTIC CELLS

  36. Electrolytic Cell a cell that uses electrical energy to produce a chemical change that would otherwise NOT occur spontaneously

  37. Process referred to as electrolysis

  38. (+) (-) M+(aq) M M X-(aq)

  39. e- e- (+) (-) M+(aq) M M X-(aq) Anode M M+ + e- oxidation Cathode M+ + e- M reduction

  40. Ampere a unit of electrical current equal to one coulomb of charge per second coul sec 1 amp = 1

  41. Coulomb a unit of electric charge equal to the quantity of charge in about 6 x 1019 electrons

  42. Faraday a constant representing the charge on one mole of electrons 1 F = 96,485 C 96,500 C

  43. 3: It is necessary to replate a silver teapot with 15.0 g of silver. If the electrolytic cell runs at 2.00 amps, how long will it take to plate the teapot?

  44. 4: Sodium metal and chlorine gas are prepared industrially in a Down’s Cell from the electrolysis of molten NaCl. What mass of metal and volume of gas can be made per day if the cell operates at 7.0 volts and 4.0 x 104 amps if the cell is 75% efficient?

  45. 5: At what current must a cell be run in order to produce 5.0 kg of aluminum in 8.0 hours if the cell produces solid aluminum from molten aluminum chloride?

  46. ELECTROCHEMISTRY, FREE ENERGY, & EQUILIBRIUM

  47. thus: wmax = - q . Emax

  48. but: wmax = G and q = nF thus if: wmax = - q . Emax then G = - nFE

  49. G = G0 + RT ln Q G = - nFE - nFE = - nFE0 + RT ln Q

  50. NERNST EQUATION

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