1. Pgs. 652 - 654

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

7.  Electronegativity = how much you “love” electrons More electronegative = more you “love” electrons = more likely to ________

8. Activity Series Let’s look at Mg and Cu: Mg + CuCl 2  Cu + MgCl 2 Cu + MgCl 2  Mg + Cu Cl 2 Li Rb K Cs Ba Sr Ca Na Mg Al Zn Cr Fe Ni Sn Pb Cu Hg Ag Au Most reactive Least reactive For a reaction to happen the solid metal must be above the aqueous metal in the activity series

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) + Cu 2+ (aq)  Fe 2+ (aq) + Cu (s)

11. Copper Plating – An Example CuSO 4 (aq) Fe FeSO 4 (aq) Zn Cu Fe Since the copper is plating the iron, the solution will get lighter as more copper is used.

12. Does the reverse happen? Fe (s) + Cu 2+ (aq)  Fe 3+ (aq) + Cu (s) Can we go backwards?…. Fe (s) + Cu 2+ (aq)  Fe 3+ (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. Cu 2+ + 2e-  Cu Fe 2+  Fe 3+ + e - Cu2+

16. Why won’t the reaction continue?? Cu 2+ + 2e-  Cu Fe 2+  Fe 3+ + e - Build up of charges would require large amounts of energy Solutions must be connected to allow ions to flow!

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

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

22. Let’s practice drawing a Cu/Zn Galvanic Cell  Cu 2+ + 2e-  CuCathode/reduction  Zn  Zn 2+ + 2e- Anode/oxidation  Cu 2+ + Zn  Zn 2+ + Cu Zn Cu SO 4 2- Zn 2+

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|Zn 2+ ||Cu 2 +|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.