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Electrochemistry(a)
- 2. Electrochemistry and Redox • Oxidation-reduction: “Redox” • Electrochemistry: • study of the interchange between chemical change and electrical work • Electrochemical cells: • systems utilizing a redox reaction to produce or use electrical energy
- 3. Redox Review • Redox reactions:electron transfer processes • Oxidation: loss of 1 or more e- • Reduction: gain of 1 or more e- • Oxidation numbers: imaginary charges • (Balancing redox reactions)
- 4. Oxidation Numbers (O.N.) • 1. Pure element O.N. is zero • 2. Monatomic ion O.N. is charge • 3. Neutral compound: sum of O.N. is zero • Polyatomic ion: sum of O.N. is ion’s charge • *Negative O.N. generally assigned to more electronegative element
- 5. Oxidation Numbers (O.N.) • 4. Hydrogen • assigned +1 • (metal hydrides, -1) • 5. Oxygen • assigned -2 • (peroxides, -1; OF2, +2) • 6. Fluorine • always -1
- 6. Oxidation-reduction • Oxidation is loss of e- • O.N. increases (more positive) • Reduction is gain of e- • O.N. decreases (more negative) • Oxidation involves loss OIL
- 7. Redox • Oxidation is loss of e- • causes reduction • “reducing agent” • Reduction is gain of e- • causes oxidation • “oxidizing agent”
- 9. Balancing Redox Reactions • 1. Write separate equations (half- reactions) for oxidation and reduction • 2. For each half-reaction • a. Balance elements involved in e- transfer • b. Balance number e- lost and gained • 3. To balance e- • multiply each half-reaction by whole numbers
- 10. Balancing Redox Reactions: Acidic • 4. Add half-reactions/cancel like terms (e- ) • 5. Acidic conditions: • Balance oxygen using H2O • Balance hydrogen using H+ • Basic conditions: • Balance oxygen using OH- • Balance hydrogen using H2O • 6. Check that all atoms and charges balance
- 11. Examples • Acidic conditions: • • Basic conditions: +++ + →+ 3 (aq) 2 (aq) 2 (aq) - 4(aq) FeMnFeMnO acid −− →++ 2(aq)2(g)(aq)(s) Ag(CN)OCNAg base
- 12. Types of cells • Voltaic (galvanic) cells: • a spontaneous reaction generates electrical energy • Electrolytic cells: • absorb free energy from an electrical source to drive a nonspontaneous reaction
- 13. Common Components • Electrodes: • conduct electricity between cell and surroundings • Electrolyte: • mixture of ions involved in reaction or carrying charge • Salt bridge: • completes circuit (provides charge balance)
- 14. Electrodes• Anode: • Oxidation occurs at the anode • Cathode: • Reduction occurs at the cathode • Active electrodes: participate in redox
- 15. Voltaic (Galvanic) Cells • A device in which chemical energy is changed to electrical energy. • Uses a spontaneous reaction.
- 16. 17_360 Porous disk Reducing agent Oxidizing agent e – e – e – e – e – e – CathodeAnode (b)(a) Oxidation Reduction
- 20. Zn2+ (aq) + Cu(s) → Cu2+ (aq) + Zn(s) • Zn gives up electrons to Cu – “pushes harder” on e- – greater potential energy – greater “electrical potential” • Spontaneous reaction due to – relative difference in metals’ abilities to give e- – ability of e- to flow
- 21. Cell Potential • Cell Potential / Electromotive Force (EMF): • The “pull” or driving force on electrons • Measured voltage (potential difference) V C J movedchargeofunit energypotentialelectricalorwork Ecell ===
- 22. 17_363 e– e– e– e– Zn 2+ SO4 2– Zn(s) 1.0 M Zn 2+ solution Anode 1.0 M Cu 2+ solution Cathode Cu 2+ SO4 2– Cu(s) Ecell = +1.10 V
- 23. Cell Potential, E0 cell • E0 cell • cell potential under standard conditions • • elements in standard states (298 K) • solutions: 1 M • gases: 1 atm
- 25. Standard Reduction Potentials• E0 values for reduction half-reactions with solutes at 1M and gases at 1 atm • Cu2+ + 2e− → Cu • E0 = 0.34 V vs. SHE • SO4 2− + 4H+ + 2e− → H2SO3 + H2O • E0 = 0.20 V vs. SHE
- 31. E0 cell and ∆G0 • E0 cell > 0 ∆G0 < 0 Spontaneous • E0 cell < 0 ∆G0 > 0 Not • E0 cell = 0 ∆G0 = 0 Equilibrium
- 32. Calculating E0 cell • E0 cell = E0 cathode - E0 anode • Br2(aq)+2V3+ +2H2O(l) → 2VO2+ (aq)+ 4H+ (aq)+ 2Br- (aq) • Given: E0 cell = +1.39 V • E0 Br2 = +1.07 V • What is E0 V3+ and is the reaction spontaneous?
- 33. E0 values • More positive: • Stronger oxidizing agent • More readily accepts e- • More negative: • Stronger reducing agent • More readily gives e- • Stronger R.A. + O.A. → Weaker R.A. + O.A.
- 34. Free Energy and Cell Potential • n: number of moles of e- • F: Faraday’s constant • 96485 C • mol of e- 00 max nFEG −=∆=w
- 35. ∆G0 , E0 , and K • At equilibrium: ∆G0 = 0 and K = Q • At 298 K: 00 nFERTlnKG −=−=∆ lnK nF RT E0 =so logK n 0.0592 E0 =
- 37. Nernst Equation• Under nonstandard conditions RTlnQnFEnFE RTlnQGG 0 0 +−=− +∆=∆ lnQ nF RT EE 0 cell −= lnQ n 0.0592 EE 0298K cell −=
- 39. Concentration Cells • . . . a cell in which both compartments have the same components but at different concentrations
- 40. 17_366 e– e–e– e– Ag 1 M Ag+ 1 M NO3 – Anode Cathode Porous disk Ag 0.1 M Ag+ 0.1 M NO3 –
- 42. 17_369 Reference solution of dilute hydrochloric acid Silver wire coated with silver chloride Thin-walled membrane
- 45. Batteries • A battery is a galvanic cell or, more commonly, a group of galvanic cells connected in series.
- 53. Fuel Cells • Galvanic cells • Reactants are continuously supplied. • 2H2(g) + O2(g) → 2H2O(l) • anode: 2H2 + 4OH− → 4H2O + 4e− • cathode: 4e− + O2 + 2H2O → 4OH−
- 56. Corrosion • Some metals, such as copper, gold, silver and platinum, are relatively difficult to oxidize. • These are often called noble metals.
- 60. Electrolysis • Forcing a current through a cell to produce a chemical change for which the cell potential is negative.
- 65. 17_378 Molten aluminum Carbon dioxide formed at the anodes Carbon-lined iron tank Plug Molten Al2O3/Na3AlF6 mixture Electrodes of graphite rods To external power source
- 66. Stoichiometry How much chemical change occurs with the flow of a given current for a specified time? • current and time → quantity of charge → moles of electrons → moles of analyte → grams of analyte