1. ELECTROCHEMISTRY Chapter 21

2. Section 21.1 Electrochemical Cells Objectives –Interpret an activity series and identify the elements that are most easily oxidized and those that are least easily oxidized –Name the type of reactions involved in electrochemical processes –Describe how a voltaic cell produces electrical energy

3. Fireflies –Glow to attract mates Anglerfish –Emit light to attract prey Squid, jellyfish, bacteria, and shrimp –Luminous How? Redox reactions Transfer of electrons in a redox reaction provides energy

4. Electrochemical Processes Chemical processes either release or absorb energy –Energy is sometimes in the form of electricity Electron transfer reactions or redox reactions –result in the generation of an electric current spontaneously or –are caused by imposing an electric current nonspontaneously –The field of chemistry that deals with these two situations is called electrochemistry

5. Spontaneous Redox Reactions How do we know this reaction between zinc and copper is spontaneous? –Table J – Activity Series Zinc is higher on the list than copper –For any two metals on table J, the more active metal is more readily oxidized

6. Which are spontaneous & if spontaneous then Identify Reaction Products ? Li + AlCl 3  Cs + CuCl 2  I 2 + NaCl  Cl 2 + KBr  Fe + CaBr 2  Mg + Sr(NO 3 ) 2  F 2 + MgCl 2  Spontaneous (Li above Al) Spontaneous (Cs / Cu) Nonspontaneous (Cl / I) Spontaneous (Cl / Br) Nonspontaneous (Ca / Fe) Nonspontaneous (Sr / Mg) Spontaneous (F / Cl)

7. Electrochemistry Many applications –Flashlights –Automobile batteries –Manufacture of sodium and aluminum metal –Silver plating of table ware and jewelry –Biological systems Carrying impulses

8. Electrochemical Cells An electrochemical cell is any device that converts chemical energy into electrical energy or electrical energy into chemical energy –Redox reactions occur in all electrochemical cells.A battery functions by transferring electrons through an external wire from the reducing agent to the oxidizing agent.

9. Overview of Electrochemistry Two kinds of electrochemical cells (kind of opposites): 1.Voltaic Use a spontaneous reaction to produce a flow of electrons (electricity) - exothermic. 2.Electrolytic Use a flow of electrons (electricity) to force a nonspontaneous reaction to occur - endothermic. Batteries are voltaic cells

10. Animation of Voltaic Cell http://www.chem.iastate.edu/group/Greenbo we/sections/projectfolder/flashfiles/electroCh em/volticCell.html http://www.chem.iastate.edu/group/Greenbo we/sections/projectfolder/flashfiles/electroCh em/volticCell.html http://www.chem.iastate.edu/group/Greenbo we/sections/projectfolder/flashfiles/electroCh em/volticCell.html http://www.chem.iastate.edu/group/Greenbo we/sections/projectfolder/flashfiles/electroCh em/volticCell.html

11. Voltaic Cells Spontaneous redox reactions Converts chemical energy into electrical energy –Electrical energy is produced in a voltaic cell by spontaneous redox reactions within the cell Parts of a Voltaic Cell –2 half cells –2 electrodes (an anode and a cathode) –Aqueous solutions –Wire –Salt bridge

12. Voltaic Voltaic Cell AnodeCathode Aqueous solution must contain ions of same metal as electrode: here ions = Zn 2+ ions. Solution might be Zn(NO 3 ) 3(aq) or ZnSO 4(aq) Aqueous solution must contain ions of same metal as electrode: here ions = Cu 2+ ions. Solution might be Cu(NO 3 ) 3(aq) or CuSO 4(aq)

13. Voltaic Voltaic Cell AnodeCathode Wire: connects the electrodes, carries electrons (electric current) Salt bridge: Allows ions to pass from one cell to another but prevents solutions from mixing completely

14. Anode and Cathode Electrodes (the metals) –Anode Electrode at which the oxidation occurs –Electrons are produced at the anode and it is labeled the negative electrode –Cathode Electrode at which reduction occurs –Electrons are consumed at the cathode as a result the cathode is labeled the positive electrode –Neither electrode is really charged All parts remain neutral Moving electrons and ions balance any charge that might build up

15. What happens at the electrodes? Anode: Zinc metal slowly dissolves (looses mass) Oxidation: Zn 0  Zn 2+ + 2e - Cathode: Copper atoms are deposited as metallic copper on top of zinc (gaining mass) Reduction: Cu 2+ +2e -  Cu 0 Zn(s) + Cu 2+ (aq)  Zn 2+ (aq) + Cu(s)

16. An OxRed Cat An Ox Ate a Red Cat Anode – Oxidation –The anode = location for the oxidation half-reaction. Reduction – Cathode –The cathode = location for the reduction half-reaction.

17. How do we determine which electrode is the anode and which electrode is the cathode? Remember… –Table J The more active metal is oxidized and is therefore the anode Anode = oxidation = electron donor –Higher metal on table J Cathode = reduction = electron acceptor –Lower metal on table J

18. Zn is above Cu on Table J Zn is the anode Cu is the cathode Electrons flow from zinc to copper (through wire) Positive ions (Zn 2+ )flow from zinc to copper (through salt bridge) Negative ions (SO 4 2- )flow from copper to zinc (through salt bridge) Zn 2+

19. e-e- e-e- e-e- e-e- e-e- e-e- AnodeCathode

20. Zn --> Zn 2+ + 2e- Cu 2+ + 2e- --> Cu <-- Anions Cations --> Oxidation Anode Negative Oxidation Anode Negative Reduction Cathode Positive Reduction Cathode Positive RED CAT AN OX Complete Electrochemical Cell

21. How a Voltaic Cell Works … The electrochemical process that occurs in a Zn-Cu voltaic cell can best be described in steps, but the steps actually occur at the same time 1.Electrons are produced at zinc anode according the oxidation half reaction: Zn (s)  Zn 2+ (aq) + 2e -. Zinc is the anode and the negative electrode because it is oxidized 2.Electrons leave zinc anode and pass through the external circuit (wire) to the copper rod 3.Electrons enter the copper rod and interact with copper ions in solution. The following reduction half reaction occurs Cu 2+ (aq) + 2e -  Cu (s). Copper ions are reduced. The copper is the cathode and the positive electrode because it is reduced 4.To complete the circuit, both the positive and negative ions move through the aqueous solution via the salt bridge

22. Construct a Voltaic Cell with Al & Pb 1.Use Table J to identify anode & cathode. 2.Draw Cell, put in electrodes & solutions using nitrate as the negative ion 3.Label: a. anode b. cathode c. direction of electron flow in wire d. direction of positive ion flow in salt bridge e. positive electrode f. negative electrode. 4.Write out half reactions 5.Write the overall reaction

23. Electron flow  Al = anode Pb = cathode wire Salt bridge Al +3 & NO 3 -1 Pb +2 & NO 3 -1 Positive ion flow  - 

24. What are the half-reactions? Al  Al +3 + 3e - Pb +2 + 2e -  Pb Al metal is the electrode – it is dissolving. Al +3 ions go into the solution. (loosing mass) Pb +2 ions are in the solution. They pick up 2 electrons at the surface of the Pb electrode & plate out. (gaining mass) Al is higher on Table J = anode = oxidation Pb is lower on Table J = cathode = reduction

25. Overall Reaction 2(Al  Al +3 + 3e - ) 3(Pb +2 + 2e -  Pb) _____________________________ 2Al + 3Pb +2  2Al +3 + 3Pb

26. Which electrode is losing mass? Which electrode is gaining mass? What’s happening to the [Al +3 ] in solution? What’s happening to the [Pb +2 ]in solution? Al Pb Increasing Decreasing

27. Animation of Voltaic Cell Let’s see the animation again: http://www.chem.iastate.edu/group/Greenbo we/sections/projectfolder/flashfiles/electroCh em/volticCell.html http://www.chem.iastate.edu/group/Greenbo we/sections/projectfolder/flashfiles/electroCh em/volticCell.html Let’s see the animation again: http://www.chem.iastate.edu/group/Greenbo we/sections/projectfolder/flashfiles/electroCh em/volticCell.html http://www.chem.iastate.edu/group/Greenbo we/sections/projectfolder/flashfiles/electroCh em/volticCell.html

28. Videos http://www.youtube.com/watch?v=kfgtU9D DvdY&feature=channel (3:29 min)http://www.youtube.com/watch?v=kfgtU9D DvdY&feature=channel

29. 21.3

30. Section 21.3 – Electrolytic Cells Objectives –Distinguish between electrolytic and voltaic cells –Describe the chemical changes that take place during electrolysis –Name three ways that electrolysis is used in metal processing

31. Electrolytic vs. Voltaic Cells Voltaic cells convert chemical energy to electrical energy during a spontaneous redox reaction Electrolytic cells use an electric current to make a nonspontaneous redox reaction go forward –This process of using electrical energy to bring about a chemical change is called electrolysis Applications of electrolysis include: –Silver plating dishes and utensils –Gold-plating jewelry –Chrome plating automobile parts

32. Voltaic vs Electrolytic Cells – Student version Answer Key

33. What’s the same? Voltaic & Electrolytic Cells What’s the same? Anode is always the site of oxidation –AN OX Cathode is always the site of reduction –RED CAT Electrons always flow from anode to cathode

34. What’s different? Voltaic & Electrolytic Cells What’s different? Presence of a battery –Voltaic – no battery –Electrolytic – has battery Which metal is the anode & which is the cathode –Voltaic – anode is more active metal  spontaneous –Electrolytic – anode is less active metal  non-spontaneous Charge on anode & charge on cathode –Voltaic – anode is + –Electrolytic – cathode is -

35. Electroplating Example of Electrolysis One metal is plated onto another Used to protect the surface of the base metal from corrosion Or to make objects, such as tableware and jewelry more attractive An electric current is used to produce a chemical reaction –The object to be plated is the cathode (negative) –The metal used for platting is the anode (positive)

36. Electroplating Cu Anode  oxidation  Cu dissolves Cu 2+ ions in solution combine with excess e- on key to form Cu coating Example – copper plating a key The key is the cathode Copper is the anode e-

37. Electroplating Usually, the object to be electroplated, such as a spoon, is cast of an inexpensive metal. It is then coated with a thin layer of a more attractive, corrosion-resistant, and expensive metal, such as silver or gold. Cathode Ag + + e-  Ag 0 Silver coats spoon anode Look at the direction of e- flow, what does that tell us about where oxidation occurs? Oxidation Oxidation… so is Ag the anode or cathode? Ag 0  Ag + + 1e-

38. Our next example is the electrolysis of molten sodium chloride –This is the only way to produce sodium metal Why do you think we need to melt the sodium chloride first? –So the ions are mobile –when you hear molten NaCl – think of ions floating around in liquid Let’s examine the electrolytic cell for molten NaCl. Electrolysis of Molten Sodium Chloride

39. +- e- AnodeCathode Inert electrode Cl - Na + Cl - + - Ox: 2Cl -  Cl 2 + 2e - Red: Na + + 1e -  Na Which way do the electrons flow through the wire? Electrons always flow from anode to cathode…which is the anode? What is the charge on the anode? What is the charge on the cathode? What electrode are the Cl- attracted to? The Na+?

40.  Chlorine ions combine to form chlorine gas  Remember this is molten NaCl, when all the Cl escapes as gas, we are left with Na +- e- AnodeCathode Inert electrode Cl - Na + Cl - + - Ox: 2Cl -  Cl 2 + 2e - Red: Na + + 1e -  Na Cl 2(g) e- e- Na 0

41. +- battery Na (l) electrode half-cell Electrolysis of Molten NaCl Na + Cl - Na + Na + + e -  Na2Cl -  Cl 2 + 2e - Cl 2 (g) escapes Observe the reactions at the electrodes NaCl (l) (-) Cl - (+)

42. In the 2 beakers a strip of Cu was placed in a solution Zn(NO 3 ) 2 or AgNO 3. Which beaker had the Zn(NO 3 ) 2 & which had AgNO 3 ? Predicting Spontaneous Redox Reactions Beaker (a) was the solution of AgNO 3 and beaker (b) was the solution of Zn(NO 3 ) 2 Because there is no outside source of electricity the only reaction that can occur is a spontaneous one – Cu is above Ag therefore spontaneous, Cu is below Zn, therefore a nonspontaneous reaction and would require an outside source of energy to force the reaction to occur

43. End of Ch 21

44. Corrosion Prevention Zinc is more easily oxidized than iron, therefore the zinc will be oxidized as opposed to the iron prevent corrosion of the iron metal

45. Application: Corrosion

46. Charging a Battery 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. This is why the ammeter in your car often goes slightly higher while your battery is charging, and then returns to normal. In your car, the battery charger is called an alternator. If you have a dead battery, it could be the battery needs to be replaced OR the alternator is not charging the battery properly.

47. Predicting Redox Reactions A + BX  B + AX A & B are metals. If metal A is above metal B in Table J, the reaction is spontaneous. X + AY  Y + AX X & Y are nonmetals. If nonmetal X is above nonmetal Y in Table J, the reaction is spontaneous.

49. +- battery e-e- e-e- NaCl (l) (-)(+) cathode anode Electrolysis of Molten NaCl Na + Cl - Na + Na + + e -  Na 2Cl -  Cl 2 + 2e - cations migrate toward (-) electrode anions migrate toward (+) electrode At the microscopic level