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Electrochemistry

Electrochemistry. Chapter 21. 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. Redox Review.

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Electrochemistry

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  1. Electrochemistry Chapter 21

  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 • Reduction involves gain RIG

  7. Redox • Oxidation is loss of e- • causes reduction • “reducing agent” • Reduction is gain of e- • causes oxidation • “oxidizing agent”

  8. 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

  9. 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

  10. Examples • Acidic conditions: • Basic conditions:

  11. 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

  12. 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)

  13. Electrodes • Anode: • Oxidation occurs at the anode • Cathode: • Reduction occurs at the cathode • Active electrodes: participate in redox • Inactive: sites of ox. and red.

  14. Voltaic (Galvanic) Cells • A device in which chemical energy is changed to electrical energy. • Uses a spontaneous reaction.

  15. Oxidation Reduction

  16. 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

  17. Cell Potential • Cell Potential/ Electromotive Force(EMF): • The “pull” or driving force on electrons • Measured voltage (potential difference)

  18. Ecell = +1.10 V

  19. Cell Potential, E0cell • E0cell • cell potential under standard conditions • elements in standard states (298 K) • solutions: 1 M • gases: 1 atm

  20. 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 • SO42 + 4H+ + 2e H2SO3 + H2O • E0 = 0.20 V vs. SHE

  21. E0celland DG0 • E0cell > 0DG0 < 0 Spontaneous • E0cell < 0DG0 > 0 Not • E0cell = 0DG0 = 0 Equilibrium

  22. Calculating E0cell • E0cell = E0cathode - E0anode • Br2(aq)+2V3+ +2H2O(l) 2VO2+(aq)+ 4H+(aq)+ 2Br-(aq) • Given: E0cell = +1.39 V • E0Br2 = +1.07 V • What is E0V3+ and is the reaction spontaneous?

  23. 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.

  24. Free Energy and Cell Potential • n: number of moles of e- • F: Faraday’s constant • 96485 C • mol of e-

  25. DG0, E0, and K • At equilibrium: DG0 = 0 and K = Q • At 298 K:

  26. Nernst Equation • Under nonstandard conditions

  27. Concentration Cells • . . . a cell in which both compartments have the same componentsbut at different concentrations

  28. Batteries • A battery is a galvanic cell or, more commonly, a group of galvanic cells connected in series.

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