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Chemistry 100 – Chapter 20

Chemistry 100 – Chapter 20. Electrochemistry. Voltaic Cells. A Schematic Galvanic Cell. e -. Porous Disk. e -. e -. Reducing Agent. Oxidizing Agent. Anode. Cathode. The Galvanic Cell Defined.

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Chemistry 100 – Chapter 20

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  1. Chemistry 100 – Chapter 20 Electrochemistry

  2. Voltaic Cells

  3. A Schematic Galvanic Cell e- Porous Disk e- e- Reducing Agent Oxidizing Agent Anode Cathode

  4. The Galvanic Cell Defined • Galvanic cells – an electrochemical cell that drives electrons through an external circuit as a result of the spontaneous redox reaction occurring inside.

  5. The Zn/Cu Galvanic Cell

  6. Voltaic Cells • We expect the Zn electrode to lose mass and the Cu electrode to gain mass. • “Rules” of voltaic cells: • At the anode electrons are products. (Oxidation) • At the cathode electrons are reactants (Reduction) • Electrons flow from the anode to the cathode.

  7. The Anode and Cathode • Galvanic cells - the anode is negative and the cathode is positive. • Electrons are made to flow through an external circuit. (Rule 3.)

  8. Cell Potentials (Electromotive Force or EMF Values) • Electromotive force (emf) - aka the cell potential • the force required to push electrons through the external circuit. • Ecell is the emf of a cell (old notation). • Now talk about the cell potential!

  9. Cell Reactions • The difference in the RHS and the LHS reaction Cu2+ (aq) + Zn (s)  Cu (s) + Zn2+ (aq) • For each half reaction, we can write the reaction quotient (see Chapter 15) as follows Cu2+ (aq) + 2 e- Cu (s) Q = 1/ [Cu2+] Zn2+ (aq) + 2 e- Zn (s) Q = 1/ [Zn2+] Overall  Qcell = [Zn2+] / [Cu2+]

  10. The Cell Potential and G • From the reaction Gibbs energy

  11. The Nernst Equation • E - standard cell potential • Cell potential under standard conditions. • [Solutes] = 1.000 mole/L • T = 298.15 K • P = 1.00 atm pressure

  12. Cell Potentials

  13. Standard Reduction Potentials • We cannot measure the potential of an individual half-cell! • We assign a particular cell as being our reference cell and then assign values to other electrodes on that basis.

  14. Cell Potentials are Intensive Properties • In the previous example, the cell potential was simply the difference between the standard potential for the Sn4+/Sn2+ reduction and the Fe3+/Fe2+ reduction. • Reason: standard cell potentials are intensive quantities.

  15. The Standard Hydrogen electrode • Eo (H+/H2) half-cell = 0.000 V e- p{H2(g)} = 1.00 atm H2 (g) [H+] = 1.00 Pt gauze

  16. A Galvanic Cell With Zinc and the Standard Hydrogen Electrode. Note - [Zn2+]= [H+] = 1.000 M

  17. The Cell Equation for the Zinc-Standard Hydrogen Electrode. • The cell reaction 2 H+ (aq) + Zn (s)  H2 (g) + Zn2+ (aq) When we measure the potential of this cell Ecell = ERHS - ELHS but ERHS = E(H+/H2) = 0.000 V  Ecell = E(Zn2+/Zn) = -0.763 V

  18. The Spontaneous Direction of a Cell reaction • Examine the magnitude the of the standard cell potential! If Eo is positive, the rG is negative! Under standard conditions, the cell will proceed spontaneously in the direction written for the cell reaction.

  19. The Composition Dependence of the Cell Potential • Nernst equation • the nonstandard cell potential (Ecell) will be a function of the concentrations of the species in the cell reaction. To calculate Ecell, we must know the cell reaction and the value of Qcell.

  20. Electrochemical Series • Look at the following series of reactions Cu2+ (aq) + 2 e- Cu (s) E(Cu2+/Cu) = 0.337 V Zn2+ (aq) + 2 e- Zn (s) E(Zn2+/Zn) = -0.763 V • Zn has a thermodynamic tendency to reduce Cu2+ (aq) Pb2+ (aq) + 2 e- Pb (s) E(Pb2+/Pb) = -0.13 V Fe2+ (aq) + 2 e- Fe (s) E(Fe2+/Fe) = -0.44 V • Fe has a thermodynamic tendency to reduce Pb2+ (aq)

  21. The larger the difference between Ered values, the larger Ecell. In a voltaic (galvanic) cell (spontaneous) Ered(cathode) is more positive than Ered(anode). Differences in Reduction Potentials

  22. Oxidizing and Reducing Agents • The more positive Ered the stronger the oxidizing agent on the left. • The more negative Ered the stronger the reducing agent on the right.

  23. Spontaneous Oxidation Processes • A species on the higher to the left of the table of standard reduction potentials will spontaneously oxidize a species that is lower to the right in the table. • Any species on the right will spontaneously reduce anything that is higher to the left in the series.

  24. Oxidizing and Reducing Agents

  25. Concentration Cells • Two identical half-cells. • RHS • AgCl (s) + e- Ag (s) + Cl- (aq, 0.10 M) • LHS • AgCl (s) + e- Ag (s) + Cl- (aq, 0.50 M) • Electrolyte concentration cell – the electrodes are identical; they simply differ in the concentration of electrolyte in the half-cells.

  26. The Nernst equation for the cell

  27. Cells at Equilibrium • When the electrochemical cell has reached equilibrium Kcell = the equilibrium constant for the cell reaction. Knowing the E° value for the cell, we can estimate the equilibrium constant for the cell reaction.

  28. Equilibrium Constants from Cell Potentials • Examine the following cell. • Half-cell reactions. • Sn4+ (aq) + 2 e- Sn2+ (aq) E(Sn4+/Sn2+) = 0.15 V • Fe3+ (aq) + e- Fe2+ (aq) E (Fe3+/Fe2+) = 0.771 V • Cell Reaction • Sn4+ (aq) + 2 Fe3+ (aq)  Sn2+ (aq) + 2 Fe2+ (aq) • Ecell = (0.771 - 0.15 V) = 0.62 V

  29. Lead-Acid Battery • A 12 V car battery - 6 cathode/anode pairs each producing 2 V. Cathode: PbO2 on a metal grid in sulfuric acid: PbO2(s) + SO42-(aq) + 4H+(aq) + 2e-  PbSO4(s) + 2H2O(l). Anode: Pb: Pb(s) + SO42-(aq)  PbSO4(s) + 2e-

  30. Lead-Acid Battery • The overall electrochemical reaction is PbO2(s) + Pb(s) + 2SO42-(aq) + 4H+(aq)  2PbSO4(s) + 2H2O(l) • for which Ecell = ERHS - ELHS = (+1.685 V) - (-0.356 V) = +2.041 V. • Wood or glass-fiber spacers are used to prevent the electrodes form touching.

  31. A Picture of a Car Battery

  32. An Alkaline Battery • Anode: Zn cap: Zn(s)  Zn2+(aq) + 2e- • Cathode: MnO2, NH4Cl and carbon paste: 2 NH4+(aq) + 2 MnO2(s) + 2e-  Mn2O3(s) + 2NH3(aq) + 2H2O(l) • Graphite rod in the center - inert cathode. • Alkaline battery, NH4Cl is replaced with KOH. • Anode: Zn powder mixed in a gel:

  33. The Alkaline Battery

  34. Fuel Cells • Direct production of electricity from fuels occurs in a fuel cell. • H2-O2 fuel cell was the primary source of electricity on Apollo moon flights. • Cathode: reduction of oxygen: 2 H2O(l) + O2(g) + 4e-  4OH-(aq) • Anode: 2H2(g) + 4OH-(aq)  4H2O(l) + 4e-

  35. Fuel Cells

  36. Corrosion of Iron • Since E(Fe2+/Fe) < E(O2/H2O) iron can be oxidized by oxygen. • Cathode • O2(g) + 4H+(aq) + 4e-  2H2O(l). • Anode • Fe(s)  Fe2+(aq) + 2e-. • Fe2+ initially formed – further oxidized to Fe3+ which forms rust, Fe2O3• xH2O(s).

  37. Rusting (Corrosion) of Iron

  38. Preventing the Corrosion of Iron • Corrosion can be prevented by coating the iron with paint or another metal. • Galvanized iron - Fe is coated with Zn. • Zn protects the iron (Zn - anode and Fe - the cathode) Zn2+(aq) +2e-  Zn(s), E(Zn2+/Zn) = -0.76 V Fe2+(aq) + 2e-  Fe(s), E(Fe2+/Fe)= -0.44 V

  39. Preventing the Corrosion of Iron

  40. Preventing the Corrosion of Iron • To protect underground pipelines, a sacrificial anode is added. • The water pipe - turned into the cathode and an active metal is used as the sacrificial anode. • Mg is used as the sacrificial anode: Mg2+(aq) +2e- Mg(s), E(Mg2+/Mg) = -2.37 V Fe2+(aq) + 2e- Fe(s), E(Fe2+/Fe) = -0.44 V

  41. Corrosion Prevention

  42. Electrolysis of Aqueous Solutions • Nonspontaneous reactions require an external current in order to force the reaction to proceed. • Electrolysis reactions are non-spontaneous. • In voltaic and electrolytic cells: • reduction occurs at the cathode, and • oxidation occurs at the anode.

  43. Voltaic vs.Electrolytic Cells • Electrolytic cells – electrons are forced to flow from the anode to cathode. • In electrolytic cells the anode is positive and the cathode is negative. (In galvanic cells the anode is negative and the cathode is positive.)

  44. Electrolysis of Aqueous Solutions

  45. Electrolysis of Molten Salts • Decomposition of molten NaCl. • Cathode: 2Na+(l) + 2e- 2Na(l) • Anode: 2Cl-(l)  Cl2(g) + 2e-. • Industrially, electrolysis is used to produce metals like Al.

  46. Electrolysis With Active Electrodes • Active electrodes: electrodes that take part in electrolysis. • Example: electrolytic plating.

  47. Electrolysis With Active Electrodes (cont’d) • Consider an active Ni electrode and another metallic electrode placed in an aqueous solution of NiSO4: • Anode: Ni(s)  Ni2+(aq) + 2e- • Cathode: Ni2+(aq) + 2e- Ni(s). • Ni plates on the inert electrode. • Electroplating is important in protecting objects from corrosion.

  48. Quantitative Aspects of Electrolysis • Consider the reduction of Cu2+ to Cu. Cu2+(aq) + 2e- Cu(s). • 2 mol of electrons  1 mol of Cu. How much material is obtained? Q = I t • current (I) • time (t) of the plating process.

  49. Gibbs Energy and Work • Gibbs energy – the maximum amount of useful work that can be obtained from a system. Note – if wmax is negative, then work is performed by the system and E is positive.

  50. Electrical Work • Eelectrolytic cell – external source of energy is required to force the reaction to proceed. • External emf must be greater than Ecell. • From physics: work has units watts. 1 W = 1 J/s.

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