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Galvanic cell

SAHAR ADHAM LECTUERE OF PHYSICAL CHEMISTRY. Galvanic cell. Comparison of Electrochemical Cells. galvanic. electrolytic. need power source. produces electrical current. two electrodes. anode (-) cathode (+). anode (+) cathode (-).

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Galvanic cell

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  1. SAHAR ADHAM LECTUERE OF PHYSICAL CHEMISTRY Galvanic cell

  2. Comparison of Electrochemical Cells galvanic electrolytic need power source produces electrical current two electrodes anode (-) cathode (+) anode (+) cathode (-) conductive medium salt bridge vessel E°cell < 0. E°cell > 0.

  3. Electron Transfer Reactions • Electron transfer reactions are oxidation-reduction or redox reactions. • Results in the generation of an electric current (electricity) or be caused by imposing an electric current. • Therefore, this field of chemistry is often called ELECTROCHEMISTRY.

  4. You can’t have one… without the other! • oxidation: loss of electrons • reduction: gain of electrons LEO the lion says GER! GER!

  5. WHY STUDY ELECTROCHEMISTRY? • Batteries • Corrosion • Industrial production of chemicalssuch as Cl2, NaOH, F2 and Al • Biological redox reactions The heme group

  6. Cu Zn Galvanic Cells • One ½ cell rxn. occurs in each compartment. • Zn Zn2+ + 2e– in the anode. • Cu2+ + 2e– Cu in cathode. • But not without a connection. Cathode=Reduction Anode=Oxidation SO42– SO42– Zn2+ Cu2+ Zn + Cu2+ Zn2+ + Cu

  7. 2e– 2e– Ion (“salt”) Bridge • But even with a connection of the electrodes, no current flows. • We need to allow neutrality in the solutions with a salt bridge to shift counterions. Cu Zn SO42– SO42– Zn2+ Cu2+ Zn + Cu2+ Zn2+ + Cu

  8. Cell Potential • Cell Potential or Electromotive Force (emf): The “pull” or driving force on the electrons.

  9. The voltage generated by the Zn/Cu galvanic cell is +1.1V under standard conditions. • Standard conditions are: • T = 25°C and P = 1 bar for gases. • Solids and liquids are pure. • Solutions are 1 M in all species. • E°cell is sum of ½ cell E° values. Standard Reduction Potentials, E°

  10. Cell potentials and Reduction potentials E°cell = E°reduced - E°oxidized E°cell = E°cathode - E°anode

  11. All ½ cells are catalogued as reduction reactions & assigned reduction potentials, E°. • The lower reduction potential ½ rxn is reversed to become the oxidation. E°oxidation = –E°reduction • That makes spontaneous E°cell > 0. • But E°red can’t be found w/o E°ox! ½ cell Reduction Potentials

  12. We need a standard electrode to make measurements against! • The Standard Hydrogen Electrode (SHE) • 2H+(aq) + 2e– H2(1 bar) E°  0 V • 1 bar H2 flows over a Pt electrode, and the full E°cell is assigned to the other electrode. E°SHE = 0 V. • E.g., standard calomel electrode: • Hg2Cl2(s) + 2e– 2 Hg(l) + Cl–E°SCE = +0.27V • a more physically convenient reference. Origin for Reduction Potentials

  13. Shorthand for a complete redox cell is of the form: • Anode | anodic soln. || cathodic soln. | Cathode • So making a cell of Cu corrosion, • Cu | Cu2+ || NO3–, NO(g), H+ |Pt • where all ions should be suffixed (aq) and both metals should have (s). Galvanic Line Notation

  14. Primary Battery: can not be recharged e.g. Mercury Battery Secondary Battery: rechargeable (storage batteries) e.g. Ni-Cad Battery Fuel Cell: reactants supplied from an external source e.g. H2/O2 fuel cells. Types ofgalvanic cells

  15. Mercury Battery Anode: Zn is reducing agent under basic conditions Cathode: HgO + H2O + 2e- ---> Hg + 2 OH- can not be recharged

  16. Ni-Cad Battery Anode (-) Cd + 2 OH- ---> Cd(OH)2 + 2e- Cathode (+) NiO(OH) + H2O + e- ---> Ni(OH)2 + OH- rechargeable

  17. Why rechargeable?. • It isbecause the products of the reaction are solids that the Ni-Cd battery can be recharged • The solid hydroxides are sticky, and remain in place. • If current is applied, the reaction can be driven backwards!

  18. HOW TO CHARGE? • When you charge a battery, you are forcing the electrons backwards (from the + to the -). To do this, you will need a higher voltage backwards than forwards.

  19. Why not rechargeable ? But in mercury battery the ZnO is not sticky, and doesn’t remain attached to the electrode. This battery is not rechargeable

  20. H2 as a Fuel Cars can use electricity generated by H2/O2 fuel cells. H2 carried in tanks or generated from hydrocarbons

  21. 17-44 Galvanic cells for which the reactants are continuously supplied.anode: 2H2 + 4OH 4H2O + 4ecathode :4e + O2 + 2H2O  4OH2H2(g) + O2(g)  2H2O(l) Fuel Cells

  22. Mercury batteries take advantage of the high density of Hg to be quite small: used in watches, hearing aids, calculators, etc. Lithium-iodine batteries are particularly small and lightweight, but also very long-lived Often used in pacemakers, where they can last for 10 years

  23. The End For more lessons, visit www.chalkbored.com

  24. TANKYOU FOR LISTENING مع تحياتى وشكرا لحضوركم المحاضرة سحر محمود ادهم V1.06

  25. Question 1 For a galvanic cell, the electrode at which reduction occurs is called the: A: Anode B: Cathode final Dr. Keck poll 5/50

  26. ANSWER For a galvanic cell, the electrode at which reduction occurs is called the: B: Cathode

  27. final Question 2 Dr. Keck For a galvanic cell, the electrode with negative polarity is called the: A: Anode B: Cathode poll 50/50

  28. ANSWER For a galvanic cell, the electrode with negative polarity is called the: A: Anode

  29. Which of the following statements is incorrect • In a galvanic cell, reduction occurs at the anode. • The cathode is labeled "+" in a voltaic cell. • Oxidation occurs at the anode in a voltaic cell. • Electrons flow from the anode to the cathode in all electrochemical cells. Question3

  30. In a galvanic cell, reduction occurs at the anode. ANSWER

  31. Consider the following notation for an electrochemical cell Zn|Zn2+ (1M)||Fe3+ (1M), Fe2+ (1M)|Pt What is the balanced equation for the cell reaction? • Zn(s) + 2Fe3+(aq) → 2Fe2+(aq) + Zn2+(aq) • Zn2+(aq) + 2Fe2+(aq) → Zn(s) + 2Fe3+(aq) • Zn(s) + 2Fe2+(aq) → 2Fe3+(aq) + Zn2+(aq) • Zn(s) + Fe3+(aq) → Fe2+(aq) + Zn2+(aq) • Zn(s) + Fe2+(aq) → Fe(s) + Zn2+(aq) Question 4

  32. Zn(s) + 2Fe3+(aq) → 2Fe2+(aq) + Zn2+(aq) ANSWER

  33. 50/50 Dr. Keck poll final Question 5 What is the oxidation state of nitrogen in HNO3? A: +3 B: +4 C: +5 D: -5

  34. ANSWER What is the oxidation state of nitrogen in HNO3? C: +5

  35. Question 6 Consider the following electrode potentials: Mg2+ + 2e– Mg E° = –2.37 V V2+ + 2e– V E° = –1.18 V Cu2+ + e– Cu+ E° = 0.15 V Which one of the reactions below will proceed spontaneously from left to right? a. Mg2+ + V  V2+ + Mg b. Mg2+ + 2Cu+ 2Cu2+ + Mg c. V2+ + 2Cu+ V +2+ Cu2+ d. V + 2Cu2+ V2+ + 2Cu+ e. none of these

  36. d. V + 2Cu2+ V2+ + 2Cu+ ANSWER

  37. Question 7 What is the oxidative state of iodine in IO3-? A: +7 B: +6 C: +5 D: +4 50/50 Dr. Keck poll final

  38. ANSWER What is the oxidative state of iodine in IO3-? C: +5

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