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Voltaic Cells

Voltaic Cells. The Nature of Electrochemical Cells. When iron metal is dipped into an aqueous solution of blue copper sulfate, the iron becomes copper plated Why? The iron loses e- to the copper What type of reaction is this???? Fe (s) + Cu 2+ (aq)  Fe 2+ (aq) + Cu (s). Fe.

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Voltaic Cells

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  1. Voltaic Cells

  2. The Nature of Electrochemical Cells • When iron metal is dipped into an aqueous solution of blue copper sulfate, the iron becomes copper plated • Why? • The iron loses e- to the copper • What type of reaction is this???? Fe(s) + Cu2+(aq) Fe2+(aq) + Cu (s)

  3. Fe FeSO4(aq) CuSO4(aq) Zn Copper Plating – An Example Fe Cu Since the copper is plating the iron, the solution will get lighter as more copper is used.

  4. Does the reverse happen? Fe(s) + Cu2+(aq) Fe3+(aq) + Cu (s) Can we go backwards?…. Fe(s) + Cu2+(aq) Fe3+(aq) + Cu (s) • Some metals are better reducing agents than others • (AKA: some metals lose e- easier than others.) • The reaction is only spontaneous one way… the reverse reaction requires an outside source of energy to work.

  5. The Nature of Electrochemical Cells • Electrochemical process: conversion between chemical and electrical energy • All involve redox reactions • To capture the electrical energy, the two half-reactions must be physically separated • Called electrochemical cells • Can create electricity or be used to create a chemical change

  6. Voltaic Cells • Invented by Alessandro Volta in 1800 • Voltaic cells: electrochemical cells used to convert chemical energy into electrical energy • Examples  flashlights or battery-powered calculators • Made of half cells • One part of the voltaic cell where oxidation or reduction is occurring

  7. Schematic for separating the oxidizing and reducing agents in a redox reaction. 8H+ + MnO4- + 5e- Mn2+ + 4H20 ; Fe2+  Fe3+ + e-

  8. Why won’t the reaction continue?? Build up of charges would require large amounts of energy Solutions must be connected to allow ions to flow! 8H+ + MnO4- + 5e- Mn2+ + 4H20 ; Fe2+  Fe3+ + e-

  9. Salt Bridge: contains a strong electrolyte held in place by gel Porous Disk: allows ion flow without mixing solutions Allows ions to pass between solutions, but doesn’t allow the solutions to mix

  10. Parts of a Voltaic Cell • Electrode: • Conductor in a circuit that carries electrons to a metal • Anode = oxidation • Negatively charged • Cathode = reduction • Positively charged

  11. Steps of a Voltaic Cell • e- created at anode • Shown in oxidation half-reaction • e- leave zinc and pass through wire • e- enter cathode and cause reduction • Shown in half-reaction • Positive and negative ions pass through salt bridge to finish the circut

  12. Oxidation half-reaction Zn(s) Zn2+(aq) + 2e- Reduction half-reaction Cu2+(aq) + 2e- Cu(s) Overall (cell) reaction Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s) A voltaic cell based on the zinc-copper reaction. Figure 21.5

  13. Schematic of a battery. Electron flow anode to cathode (- to +) oxidation to reduction reducing agent to oxidizing agent

  14. Cu/Zn Voltaic Cell • Cu2+ + 2e-  Cu Cathode/reduction • Zn  Zn2+ + 2e- Anode/oxidation • Cu2+ + Zn  Zn2+ + Cu Cu Zn SO42- Zn2+ Cu2+ Zn2+

  15. Voltaic Cell Shorthand • Oxidation half cell is listed first with reduced and oxidized species separated by a line. • Reduction is next in the opposite order. • Double line separates the two and represents a salt bridge. Zn|Zn2+||Cu2+|Cu

  16. Voltaic Cell Shorthand • Draw shorthand notation for a Mg-Pb cell where the nitrate ion is present. You might want to refer to the activity series to determine what is oxidized and what is reduced! • Draw a diagram for this voltaic cell on the back or bottom of your notes! • Label: anode, cathode, direction of e- flow • Write-out the ½ rxns and combined reaction.

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