1. 21.1 Electrochemical Cells > 1 BATLAB16.pdf BATLAB16.pdf Day 2: Lab part 1 BATLAB16.pdf BATLAB16.pdf 36 The Activity Series-S.pdf Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

2. 21.1 Electrochemical Cells > 2 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Cu Mg Zn Zn 2+ Mg 2+ Cu 2+ solids (aq) temp

3. 21.1 Electrochemical Cells > 3 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Electrochemical Processes Chemical processes can either release energy or absorb energy. The energy can sometimes be in the form of electricity. An electrochemical process is any conversion between chemical energy and electrical energy.

4. 21.1 Electrochemical Cells > 4 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Electrochemical Processes What type of chemical reaction is involved in all electrochemical processes? All electrochemical processes involve redox reactions.

5. 21.1 Electrochemical Cells > 5 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Electrochemical Processes When a zinc strip is dipped into a copper(II) sulfate solution, electrons are transferred from zinc atoms to copper ions. This flow of electrons is an electric current.

6. 21.1 Electrochemical Cells > 6 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Electrons are transferred from zinc atoms to copper atoms. Zn(s) + Cu 2+ (aq) → Zn 2+ (aq) + Cu(s)

7. 21.1 Electrochemical Cells > 7 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. The net ionic equation involves only zinc and copper. Zinc metal oxidizes spontaneously in a copper-ion solution. Zn(s) + Cu 2+ (aq) → Zn 2+ (aq) + Cu(s) Oxidation:Zn(s) → Zn 2+ (aq) + 2e Reduction: Cu 2+ (aq) + 2e – → Cu(s)

8. 8 CHEMICAL CHANGE ---> ELECTRIC CURRENT With time, Cu plates out onto Zn metal strip, and Zn strip “disappears.” Zn is oxidized Zn(s) ---> Zn 2+ (aq) + 2e- Cu 2+ is reduced Cu 2+ (aq) + 2e- ---> Cu(s)

9. 21.1 Electrochemical Cells > 9 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Electrochemical Processes For any two metals in an activity series, the more active metal is the more readily oxidized. Activity Series of Metals ElementOxidation half-reactions Most active and most easily oxidized LithiumLi(s) → Li + (aq) + e – BariumBa(s) → Ba 2+ (aq) + 2e – CalciumCa(s) → Ca 2+ (aq) + 2e – AluminumAl(s) → Al 3+ (aq) + 3e – ZincZn(s) → Zn 2+ (aq) + 2e – IronFe(s) → Fe 2+ (aq) + 2e – NickelNi(s) → Ni 2+ (aq) + 2e – TinSn(s) → Sn 2+ (aq) + 2e – LeadPb(s) → Pb 2+ (aq) + 2e – Hydrogen * H 2 (g) → 2H + (aq) + 2e – Least easily oxidized CopperCu(s) → Cu 2+ (aq) + 2e – SilverAg(s) → Ag + (aq) + e – MercuryHg(s) → Hg 2+ (aq) + 2e – * Hydrogen is included for reference purposes. Decreasing activity

10. 21.1 Electrochemical Cells > 10 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Electrochemical Processes Activity Series of Metals ElementOxidation half-reactions Most active and most easily oxidized LithiumLi(s) → Li + (aq) + e – BariumBa(s) → Ba 2+ (aq) + 2e – CalciumCa(s) → Ca 2+ (aq) + 2e – AluminumAl(s) → Al 3+ (aq) + 3e – ZincZn(s) → Zn 2+ (aq) + 2e – IronFe(s) → Fe 2+ (aq) + 2e – NickelNi(s) → Ni 2+ (aq) + 2e – TinSn(s) → Sn 2+ (aq) + 2e – LeadPb(s) → Pb 2+ (aq) + 2e – Hydrogen * H 2 (g) → 2H + (aq) + 2e – Least easily oxidized CopperCu(s) → Cu 2+ (aq) + 2e – SilverAg(s) → Ag + (aq) + e – MercuryHg(s) → Hg 2+ (aq) + 2e – * Hydrogen is included for reference purposes. Decreasing activity Zinc is above copper on the list. Zinc is more readily oxidized than copper. When zinc is dipped into a copper(II) sulfate solution, zinc becomes plated with copper.

11. 21.1 Electrochemical Cells > 11 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Electrochemical Processes Activity Series of Metals ElementOxidation half-reactions Most active and most easily oxidized LithiumLi(s) → Li + (aq) + e – BariumBa(s) → Ba 2+ (aq) + 2e – CalciumCa(s) → Ca 2+ (aq) + 2e – AluminumAl(s) → Al 3+ (aq) + 3e – ZincZn(s) → Zn 2+ (aq) + 2e – IronFe(s) → Fe 2+ (aq) + 2e – NickelNi(s) → Ni 2+ (aq) + 2e – TinSn(s) → Sn 2+ (aq) + 2e – LeadPb(s) → Pb 2+ (aq) + 2e – Hydrogen * H 2 (g) → 2H + (aq) + 2e – Least easily oxidized CopperCu(s) → Cu 2+ (aq) + 2e – SilverAg(s) → Ag + (aq) + e – MercuryHg(s) → Hg 2+ (aq) + 2e – * Hydrogen is included for reference purposes. Decreasing activity Based on your data, where does Mg fall on this chart?

12. 21.1 Electrochemical Cells > 12 Why is Mg now at the bottom? REDUCTION HALF- REACTIONS

13. 21.1 Electrochemical Cells > 13 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Zn is above Pb in the activity series of metals. Which of the following statements is correct? A.Zn will react with Pb 2+. B.Pb 2+ will react with Zn 2+. C.Zn 2+ will react with Pb. D.Pb will react with Zn 2+. A.Zn will react with Pb 2+.

14. 14 Electrons are transferred from Al to Cu 2+, but there is no useful electric current. Energy released as HEAT. Electrochemical Reactions

15. 15 If Al and Cu are separated an electric current is generated and work is done by the electrons. Voltmeter is used to measure the energy.

16. 21.1 Electrochemical Cells > 16 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Voltaic Cells How does a voltaic cell produce electrical ENERGY? A voltaic cell is an electrochemical cell used to convert chemical energy into electrical energy.

17. 21.1 Electrochemical Cells > 17 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Voltaic Cells A voltaic cell consists of two half-cells. A half-cell is one part of a voltaic cell in which either oxidation or reduction occurs. –A typical half-cell consists of a piece of metal immersed in a solution of its ions.

18. 18 To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs thru an external wire. ELECTRIC CURRENT This is accomplished in a GALVANIC or VOLTAIC cell.. A group of such cells is called a battery.

19. 19 AnodeCathode Basic Concepts of Electrochemical Cells

20. 21.1 Electrochemical Cells > 20 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Constructing a Voltaic Cell The electrode at which oxidation occurs is called the anode. An electrode is a conductor in a circuit that carries electrons to or from a substance other than a metal. Electrons are produced at the anode: out The anode is labeled the negative electrode in a voltaic cell.

21. 21.1 Electrochemical Cells > 21 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Constructing a Voltaic Cell The electrode at which reduction occurs is called the cathode. Electrons are consumed at the cathode: reactant The ca t hode is labeled the positive + electrode in a voltaic cell.

22. 22 Anode, site of oxidation, negative Cathode, site of reduction, positive

23. 23 Electrons travel thru external wire. Salt bridge allows anions and cations to move between electrode compartments. Electrons travel thru external wire. Salt bridge allows anions and cations to move between electrode compartments. Zn --> Zn 2+ + 2e- Cu 2+ + 2e- --> Cu <--AnionsCations--> OxidationAnodeNegativeOxidationAnodeNegative ReductionCathodePositiveReductionCathodePositive

24. 21.1 Electrochemical Cells > 24 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Voltaic Cells The half-cells are connected by a salt bridge, which is a tube containing a strong electrolyte, often potassium sulfate (K 2 SO 4 ). The salt bridge or porous plate allows ions to pass from one half-cell to the other but prevents the solutions from mixing completely. A porous plate may be used instead of a salt bridge.

25. 21.1 Electrochemical Cells > 25 BATLAB16.pdf BATLAB16.pdf Day 2: Lab part 2 BATLAB16.pdf BATLAB16.pdf Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 37 Batteries-S.pdf

26. 26 RED = REDUCED BLACK = OXIDIZED e-

27. 21.1 Electrochemical Cells > 27 Anode (–) Cathode (+) Salt bridge Cotton plugs ZnSO 4 solution CuSO 2 solution Zn(s) Zn 2+ (aq) + 2e – Cu 2+ (aq) + 2e – Cu(s) Wiree–e– e–e– Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. A Voltaic Cell In this voltaic cell, the electrons generated from the oxidation of Zn to Zn 2+ flow through the external circuit (the wire) into the copper strip.

28. 21.1 Electrochemical Cells > 28 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Voltaic Cells The zinc and copper strips in this voltaic cell serve as the electrodes. Constructing a Voltaic Cell Anode (–) Cathode (+) Salt bridge Cotton plugs ZnSO 4 solution CuSO 2 solution Zn(s) Zn 2+ (aq) + 2e – Cu 2+ (aq) + 2e – Cu(s) Wiree–e– e–e–

29. 21.1 Electrochemical Cells > 29 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Voltaic Cells How a Voltaic Cell Works Step 1 Electrons are produced at the zinc strip according to the oxidation half-reaction: Zn(s) → Zn 2+ (aq) + 2e – Anode (–) Cathode (+) Salt bridge Cotton plugs ZnSO 4 solution CuSO 2 solution Zn(s) Zn 2+ (aq) + 2e – Cu 2+ (aq) + 2e – Cu(s) Wiree–e– e–e–

30. 21.1 Electrochemical Cells > 30 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Voltaic Cells How a Voltaic Cell Works The electrons leave the zinc anode and pass through the external circuit to the copper strip. Step 2 Anode (–) Cathode (+) Salt bridge Cotton plugs ZnSO 4 solution CuSO 2 solution Zn(s) Zn 2+ (aq) + 2e – Cu 2+ (aq) + 2e – Cu(s) Wiree–e– e–e– If a lightbulb is in the circuit, the electron flow will cause it to light.

31. 21.1 Electrochemical Cells > 31 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Voltaic Cells How a Voltaic Cell Works Step 3 Electrons interact with copper ions in solution. There, the following reduction half-reaction occurs: Cu 2+ (aq) + 2e – → Cu(s) Anode (–) Cathode (+) Salt bridge Cotton plugs ZnSO 4 solution CuSO 2 solution Zn(s) Zn 2+ (aq) + 2e – Cu 2+ (aq) + 2e – Cu(s) Wiree–e– e–e–

32. 21.1 Electrochemical Cells > 32 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Voltaic Cells How a Voltaic Cell Works Step 4 To complete the circuit, both positive and negative ions move through the aqueous solutions via the salt bridge. Anode (–) Cathode (+) Salt bridge Cotton plugs ZnSO 4 solution CuSO 2 solution Zn(s) Zn 2+ (aq) + 2e – Cu 2+ (aq) + 2e – Cu(s) Wiree–e– e–e–

33. 21.1 Electrochemical Cells > 33 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. How a Voltaic Cell Works The two half-reactions can be summed to show the overall reaction. Note that the electrons must cancel. Zn(s) → Zn 2+ (aq) + 2e – Cu 2+ (aq) + 2e – → Cu(s) Zn(s) + Cu 2+ (aq) → Zn 2+ (aq) + Cu(s)

34. 21.1 Electrochemical Cells > 34 The half-cell that undergoes oxidation (the anode) is written first, to the left of the double vertical lines. Representing Electrochemical Cells The single vertical lines indicate boundaries of phases that are in contact. The double vertical lines represent the salt bridge or porous partition that separates the anode compartment from the cathode compartment. You can represent the zinc-copper voltaic cell by using the following shorthand form. Zn(s) ZnSO 4 (aq) CuSO 4 (aq) Cu(s)

35. 35 Terms Used for Voltaic Cells Figure 20.3

36. 36 The Cu|Cu 2+ and Ag|Ag + Cell

37. 37

38. 38 CATHODEANODE Mg(s) → Mg 2+ (aq) + 2e – 2Ag 1+ (aq) + 2e – → 2Ag(s) Mg(s) + 2Ag + (aq) → Mg 2+ (aq) + 2Ag(s)

39. 21.1 Electrochemical Cells > 39 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. A voltaic cell is formed from a piece of iron in a solution of Fe(NO 3 ) 2 and silver in a solution of AgNO 3. Which is the cathode, and which is the anode? Why?

40. 21.1 Electrochemical Cells > 40 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. A voltaic cell is formed from a piece of iron in a solution of Fe(NO 3 ) 2 and silver in a solution of AgNO 3. What is the cathode half-reaction? and the anode? Cell notation? Fe(s) FeNO 3 (aq) AgNO 3 (aq) Ag(s) Fe(s) → Fe 2+ (aq) + 2e – 2Ag 1+ (aq) + 2e – → 2Ag(s) Fe(s) + 2Ag + (aq) → Fe 2+ (aq) +2 Ag(s)

41. 41 Electrochemical Electrochemical Cell Electrons move from anode to cathode in the wire. Anions & cations move thru the salt bridge.

42. 21.1 Electrochemical Cells > 42 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Current Applications Electrochemical processes produce electrical energy in dry cells, lead storage batteries, and fuel cells. A dry cell is a voltaic cell in which the electrolyte is a paste. A battery is a group of voltaic cells connected together. Fuel cells are voltaic cells in which a fuel undergoes oxidation and from which electrical energy is continuously obtained.

43. 21.1 Electrochemical Cells > 43 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Dry Cells In one type of dry cell, a zinc container is filled with a thick, moist electrolyte paste of manganese(IV) oxide (MnO 2 ), zinc chloride (ZnCl 2 ), ammonium chloride (NH 4 Cl), and water (H 2 O). Positive button (+) Graphite rod (cathode) Moist paste of MnO 2, ZnCl 2, NH 4 Cl 2, H 2 O, and graphite powder Zinc (anode) Negative end cap (–) A graphite rod is embedded in the paste. The zinc container is the anode, and the graphite rod is the cathode. The thick paste and its surrounding paper liner prevent the contents of the cell from freely mixing, so a salt bridge is not needed.

44. 21.1 Electrochemical Cells > 44 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Dry Cells Oxidation: Zn(s) → Zn 2+ (aq) + 2e – (at anode) Reduction: 2MnO 2 (s) + 2NH 4 + (aq) + 2e – → Mn 2 O 3 (s) + 2NH 3 (aq) + H 2 O(l) (at cathode) Dry cells of this type are not rechargeable because the cathode reaction is not reversible. Positive button (+) Graphite rod (cathode) Moist paste of MnO 2, ZnCl 2, NH 4 Cl 2, H 2 O, and graphite powder Zinc (anode) Negative end cap (–)

45. 45 Nearly same reactions as in common dry cell, but under basic conditions. Alkaline Battery Anode (-): Zn + 2 OH - ---> ZnO + H 2 O + 2e- Cathode (+): 2 MnO 2 + H 2 O + 2e- ---> Mn 2 O 3 + 2 OH - Negative end cap (–) Zinc (anode) Absorbent separator Graphite rod (cathode) MnO 2 in KOH paste Positive button (+) Steel case

46. 21.1 Electrochemical Cells > 46 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Lead storage batteries A battery is a group of voltaic cells connected together. A 12-V car battery consists of six voltaic cells connected together.

47. 21.1 Electrochemical Cells > 47 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. The half-reactions are as follows: Oxidation: Pb(s) + SO 4 2– (aq) → PbSO 4 (s) + 2e – Reduction: PbO 2 (s) + 4H + (aq) + SO 4 2– (aq) + 2e – → PbSO 4 (s) + 2H 2 O(l) Lead Storage Batteries

48. 21.1 Electrochemical Cells > 48 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Lead Storage Batteries The overall spontaneous redox reaction that occurs is the sum of the oxidation and reduction half-reactions. Pb + PbO 2 + 2H 2 SO 4 (aq) → 2PbSO 4 + 2H 2 O Lead(II) sulfate forms during discharge.The sulfate slowly builds up on the plates, and the concentration of the sulfuric acid electrolyte decreases.

49. 21.1 Electrochemical Cells > 49 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Lead Storage Batteries The reverse reaction occurs when a battery is recharged. 2PbSO 4 + 2H 2 O → Pb + PbO 2 + 2H 2 SO 4 (aq) This is not a spontaneous reaction. To make this reaction proceed, a direct current must be applied and pass through the cell

50. 50 Ni-Cad Battery Anode (-) Cd + 2 OH - ---> Cd(OH) 2 + 2e- Cathode (+) NiO(OH) + H 2 O + e- ---> Ni(OH) 2 + OH -

51. 21.1 Electrochemical Cells > 51 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. The hydrogen-oxygen fuel cell is a clean source of power. The only product of the reaction is liquid water. Such cells can be used to fuel vehicles. Fuel Cells

52. 21.1 Electrochemical Cells > 52 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. FUEL CELLS The half-reactions are as follows: Oxidation: 2H 2 (g) → 4H + (aq) + 4e – (at anode) Reduction: O 2 (g) + 4H + (aq) + 4e – → 2H 2 O(g) (at cathode) The overall reaction is the oxidation of hydrogen (fuel) to form water. 2H 2 (g) + O 2 (g) → 2H 2 O(g)

53. 21.1 Electrochemical Cells > 53 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Fuel Cells Fuel cells are voltaic cells in which a fuel substance undergoes oxidation and from which electrical energy is continuously obtained. They can be designed to emit no air pollutants and to operate more quietly and more cost-effectively than a conventional electrical generator.