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Utilizes relationship between chemical potential energy & electrical energy

Explore the relationship between chemical potential energy and electrical energy in electrochemistry. Learn about redox reactions, battery usage, corrosion prevention, and more. Predict redox reactions and understand the components of electrochemical cells.

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Utilizes relationship between chemical potential energy & electrical energy

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  1. Electrochemistry Utilizes relationship between chemical potential energy & electrical energy

  2. Redox Reactions • Need battery to start car • Prevent corrosion • Bleach is an oxidizing agent • Na, Al, Cl prepared or purified by redox reactions • Breathing • O2 H2O and CO2

  3. Redox Reactions • Synthesis • Decomposition • Single Replacement • Double Replacement only is not redox Redox

  4. Predicting Redox Reactions • Use Table J to predict if a given redox reaction will occur. • Any metal will donate its electrons to the ion of any metal below it. • Any nonmetal will steal electrons from the ion of any nonmetal below it. Memory Jogger

  5. Predicting Single Replacement Redox Reactions • Element + Compound  New Element + New Compound • If the element is above the swapable ion, the reaction is spontaneous. • If the element is below the swapable ion, the reaction is not spontaneous. Memory Jogger

  6. 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. Memory Jogger

  7. Which are spontaneous? Yes • Li + AlCl3 • Cs + CuCl2  • I2 + NaCl  • Cl2 + KBr  • Fe + CaBr2  • Mg + Sr(NO3)2  • F2 + MgCl2  Yes No Yes No No Yes

  8. Started with Zn(NO3)2 & Cu and AgNO3 & Cu. Which beaker had the Zn ions & which had the Ag ions?

  9. Overview of Electrochemistry • TWO kinds of cells (kind of opposites): • Galvanic or Voltaic (NYS – Electrochemical) • Use a spontaneous reaction to produce a flow of electrons (electricity). Exothermic. • Electrolytic • Use a flow of electrons (electricity) to force a nonspontaneous reaction to occur. Endothermic.

  10. Redox Half-reaction Oxidation Reduction Cell Half-Cell Electrode Anode Cathode Galvanic Voltaic Electrochemical Electrolytic Salt bridge Vocabulary

  11. Electrochemical Cells • Use a spontaneous single replacement redox reaction to produce a flow of electrons. • Electrons flow from oxidized substance to reduced substance. • Called: Galvanic cells, voltaic cells, or electrochemical cells (NYS)

  12. Electrochemical Cells • Redox reaction is arranged so the electrons are forced to flow through a wire. • When the electrons travel through a wire, we can make them do work, like light a bulb or ring a buzzer. • So the oxidation & reduction reactions have to be separated physically. OJ clock

  13. Al / CuCl2 Lab • Was a redox reaction. • Did NOT force electrons to travel through a wire. Got NO useful work out of system. • Have to be clever in how we arrange things.

  14. 2Al + 3Cu+2 2Al+3 + 3Cu Got no useful work because half-reactions weren’t separated.

  15. Half-Cell • Where each of the half-reactions takes place. • Need 2 half-cells to have a complete redox reaction. • Need to be connected by a wire for the electrons to flow through. • Need to be connected by a salt bridge to maintain electrical neutrality.

  16. Schematic of Galvanic Cell

  17. Parts of a Voltaic Cell • 2 half-cells: oxidation & reduction • Each half-cell consists of a container of an aqueous solution & an electrode or surface at which the electron transfer takes place. • Wire connecting electrodes. • Salt bridge connects solutions.

  18. How much work can you get out of this reaction? • You can measure the voltage by making the electrons travel through a voltmeter. • The galvanic cell is a battery. Of course, it’s not a very easy battery to transport or use in real-life applications.

  19. Electrode Surface at which oxidation or reduction half-reaction occurs. Anode & Cathode

  20. 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.

  21. Anode / Cathode • How do you know which electrode is which? • Use Table J to predict which electrode is the anode and which electrode is the cathode.

  22. Anode • Anode = Oxidation = Electron Donor • The anode is the metal that’s higher in Table J.

  23. Cathode • Cathode = Reduction = Electron Acceptor • The cathode is the metal that’s lower in Table J.

  24. Zn is above Cu, Zn is anode http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/galvan5.swf

  25. Notation for Cells ZnZn+2Cu+2Cu

  26. Direction of Electron Flow(wire) Anode to Cathode Direction of Positive Ion Flow (salt bridge) Anode to Cathode

  27. Positive & Negative Electrode • Negative electrode is where electrons originate – here it’s the Zn electrode. • Positive electrode is electrode that attracts electrons – here it’s the Cu electrode.

  28. Aqueous Solution • Solution containing ions of the same element as the electrode. • Cu electrode: solution may be Cu(NO3)3 or CuSO4. • Zn electrode: solution may be Zn(NO3)2 or ZnSO4.

  29. Salt Bridge • Allows for migration of ions between half-cells. • Necessary to maintain electrical neutrality. • Reaction will not proceed without salt bridge.

  30. A(s) + BX(aq)  B(s) + AX(aq) • Single replacement rxn occurs during operation of galvanic cell. • One electrode will gain mass (B) and one electrode will dissolve (A). • The concentration of metal ions will increase in one solution (making AX) & decrease in one solution (using up BX).

  31. Half-Reactions Zn  Zn+2 + 2e- Cu+2 + 2e-  Cu _________________________ Zn + Cu+2 Zn+2 + Cu Which electrode is dissolving? Which species is getting more concentrated? Zn Zn+2

  32. Zn + Cu+2 Zn+2 + Cu Cu • Which electrode is gaining mass? • Which species is getting more dilute? Cu+2

  33. When the reaction reaches equilibrium • The voltage goes to 0.

  34. Construct Galvanic Cell with Al & Pb • Use Table J to identify anode & cathode. • Draw Cell, put in electrodes & solutions • Label anode, cathode, direction of electron flow in wire, direction of positive ion flow in salt bridge, positive electrode, negative electrode. • Negative electrode is where electrons originate. Positive electrode attracts electrons.

  35. Electron flow  wire Positive ion flow  Pb = cathode Al = anode Salt bridge -  Pb+2 & NO3-1 Al+3 & NO3-1

  36. What are half-reactions? Al  Al+3 + 3e- Pb+2 + 2e-  Pb Al metal is the electrode – it’s dissolving. Al+3 ions go into the solution. Pb+2 ions are in the solution. They pick up 2 electrons at the surface of the Pb electrode & plate out.

  37. Overall Rxn 2(Al  Al+3 + 3e-) 3(Pb+2 + 2e-  Pb) _____________________________ 2Al + 3Pb+2 2Al+3 + 3Pb

  38. 2Al + 3Pb+2 2Al+3 + 3Pb Al • Which electrode is losing mass? • Which electrode is gaining mass? • What’s happening to the [Al+3]? • What’s happening to the [Pb+2]? Pb Increasing Decreasing

  39. Application: Batteries

  40. Dry Cell

  41. Mercury battery

  42. Fuel Cell • Converts chemical to electrical energy • Use oxygen or other oxidizing agent • Constant supply of fuel (hydrogen or other hydrocarbon)

  43. Application: Corrosion

  44. Corrosion Prevention

  45. What’s wrong with this picture?

  46. Daniell Cell • Invented in 1836 by John Frederic Daniell • Improved battery technology (voltaic pile –problem with hydrogen bubbling) • Definition of the volt

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