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Electrochemistry

Pgs. 652 - 654. Electrochemistry. How does our lab from Friday link to corrosion?. Corrosion is the process of returning metals to their natural state It’s a REDOX reaction!! Fe (s) + O2 (g)  Fe2O3 (s). LOTS of metals corrode, but not all of them corrode to the same extent:

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Electrochemistry

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  1. Pgs. 652 - 654 Electrochemistry

  2. How does our lab from Friday link to corrosion? • Corrosion is the process of returning metals to their natural state • It’s a REDOX reaction!! Fe (s) + O2 (g)  Fe2O3 (s)

  3. LOTS of metals corrode, but not all of them corrode to the same extent: • Ex  Aluminum!! • Aluminum will be oxidized by the air • Al (s) + O2 (g)  Al2O3 (s) • A thin layer of Al2O3 will cover the metal and protect it from further corrosion

  4. How can we protect these metals from corrosion? • The Mg will react instead of the iron…but why???

  5. So what does this have to do with the lab?? • It all comes down to HOW ACTIVE a metal is!! • What was the most active metal you saw in the lab? • What was the least active metal? • What does it all mean???

  6. Most reactive Activity Series How does electronegativity relate? Least reactive Li Rb K CsBa Sr Ca Na Mg Al Zn Cr Fe Ni Sn Pb Cu Hg Ag Au

  7. How does electronegativity relate? More electronegative = more you “love” electrons = more likely to ________ Electronegativity = how much you “love” electrons

  8. Most reactive Activity Series Let’s look at Mg and Cu: Mg + CuCl2 Cu + MgCl2 Cu + MgCl2  Mg + Cu Cl2 For a reaction to happen the solid metal must be above the aqueous metal in the activity series Least reactive Li Rb K CsBa Sr Ca Na Mg Al Zn Cr Fe Ni Sn Pb Cu Hg Ag Au

  9. Electrochemistry • The study of the interchange of chemical and electrical energy • Two types of processes in electrochemistry: • The production of an electric current from a chemical (redox) reaction • The use of an electric current to produce a chemical change

  10. But first, a demo review from yesterday… • 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)

  11. 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.

  12. 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.

  13. What does all of this mean? • 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!!

  14. Galvanic Cells • Invented by Alessandro Volta in 1800 • Galvanic cells: electrochemical cells used to convert chemical energy into electrical energy • Examples  alkaline batteries • Made of half cells • One part of the galvanic cell where oxidation or reduction is occurring

  15. Schematic for separating the oxidizing and reducing agents in a redox reaction. Cu2+ Cu2+ + 2e-  Cu Fe2+  Fe3+ + e-

  16. 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! Cu2+ + 2e-  Cu Fe2+  Fe3+ + e-

  17. 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

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

  19. Steps of a Galvanic 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 circuit

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

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

  22. Let’s practice drawing a Cu/Zn Galvanic Cell • Cu2+ + 2e-  Cu Cathode/reduction • Zn  Zn2+ + 2e- Anode/oxidation • Cu2+ + Zn  Zn2+ + Cu Cu Zn SO42- Zn2+

  23. 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 and electron transfer: Zn|Zn2+||Cu2+|Cu

  24. 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 galvanic cell on the scratch paper provided! • Label: anode, cathode, direction of e- flow • Write-out the ½ rxns and combined reaction.

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