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Chapter 20 Principles of Reactivity: Electron Transfer Reactions

Chapter 20 Principles of Reactivity: Electron Transfer Reactions. Important – Read Before Using Slides in Class

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Chapter 20 Principles of Reactivity: Electron Transfer Reactions

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  1. Chapter 20Principles of Reactivity: Electron Transfer Reactions

  2. Important – Read Before Using Slides in Class Instructor: This PowerPoint presentation contains photos and figures from the text, as well as selected animations and videos. For animations and videos to run properly, we recommend that you run this PowerPoint presentation from the PowerLecture disc inserted in your computer. Also, for the mathematical symbols to display properly, you must install the supplied font called “Symb_chm,” supplied as a cross-platform TrueType font in the “Font_for_Lectures” folder in the "Media" folder on this disc. If you prefer to customize the presentation or run it without the PowerLecture disc inserted, the animations and videos will only run properly if you also copy the associated animation and video files for each chapter onto your computer. Follow these steps: 1. Go to the disc drive directory containing the PowerLecture disc, and then to the “Media” folder, and then to the “PowerPoint_Lectures” folder. 2. In the “PowerPoint_Lectures” folder, copy the entire chapter folder to your computer. Chapter folders are named “chapter1”, “chapter2”, etc. Each chapter folder contains the PowerPoint Lecture file as well as the animation and video files. For assistance with installing the fonts or copying the animations and video files, please visit our Technical Support at http://academic.cengage.com/support or call (800) 423-0563. Thank you.

  3. ELECTROCHEMISTRYChapter 20

  4. TRANSFER REACTIONS Atom/Group transfer HCl + H2O f Cl- + H3O+ Electron transfer Cu(s) + 2 Ag+(aq) f Cu2+(aq) + 2 Ag(s)

  5. Electron Transfer Reactions • Electron transfer reactions are oxidation-reduction or redox reactions. • Redox reactions can result in the generation of an electric current or be caused by imposing an electric current. • Therefore, this field of chemistry is often called ELECTROCHEMISTRY.

  6. Review of Terminology for Redox Reactions • OXIDATION—loss of electron(s) by a species; increase in oxidation number. • REDUCTION—gain of electron(s); decrease in oxidation number. • OXIDIZING AGENT—electron acceptor; species is reduced. • REDUCING AGENT—electron donor; species is oxidized.

  7. OXIDATION-REDUCTION REACTIONS Direct Redox Reaction Oxidizing and reducing agents in direct contact. Cu(s) + 2 Ag+(aq) f Cu2+(aq) + 2 Ag(s) PLAY MOVIE

  8. OXIDATION-REDUCTION REACTIONS Indirect Redox Reaction A battery functions by transferring electrons through an external wire from the reducing agent to the oxidizing agent. PLAY MOVIE

  9. Why Study Electrochemistry? • Batteries • Corrosion • Industrial production of chemicalssuch as Cl2, NaOH, F2 and Al • Biological redox reactions The heme group

  10. Electrochemical Cells • An apparatus that allows a redox reaction to occur by transferring electrons through an external connector. • Product favored reactionfvoltaic or galvanic cellf chemical change produces electric current • Reactant favored reaction felectrolytic cellf electric current used to cause chemical change. Batteries are voltaic cells

  11. Electrochemistry Alessandro Volta, 1745-1827, Italian scientist and inventor. Luigi Galvani, 1737-1798, Italian scientist and inventor.

  12. Balancing Equations for Redox Reactions Some redox reactions have equations that must be balanced by special techniques PLAY MOVIE MnO4- + 5 Fe2+ + 8 H+fMn2+ + 5 Fe3+ + 4 H2O Mn = +7 Fe = +2 Mn = +2 Fe = +3

  13. Balancing Equations Cu + Ag+ --give--> Cu2+ + Ag

  14. Balancing Equations Step 1: Divide the reaction into half-reactions, one for oxidation and the other for reduction. Ox Cu f Cu2+ Red Ag+f Ag Step 2: Balance each for mass. Already done in this case. Step 3: Balance each half-reaction for charge by adding electrons. Ox Cu f Cu2+ + 2e- Red Ag+ + e- f Ag

  15. Balancing Equations Step 4: Multiply each half-reaction by a factor so that the reducing agent supplies as many electrons as the oxidizing agent requires. Reducing agent Cu f Cu2+ + 2e- Oxidizing agent 2 Ag+ + 2 e- f2 Ag Step 5: Add half-reactions to give the overall equation. Cu + 2 Ag+f Cu2+ + 2Ag The equation is now balanced for both charge and mass.

  16. Reduction of VO2+ with Zn

  17. Balancing Equations Balance the following in acid solution— VO2+ + Zn f VO2+ + Zn2+ Step 1: Write the half-reactions Ox Zn f Zn2+ Red VO2+f VO2+ Step 2: Balance each half-reaction for mass. Ox Zn f Zn2+ Red 2 H++ VO2+f VO2+ + H2O Add H2O on O-deficient side and add H+ on other side for H-balance.

  18. Balancing Equations Step 3: Balance half-reactions for charge. Ox Zn f Zn2+ +2e- Red e-+ 2 H+ + VO2+f VO2+ + H2O Step 4: Multiply by an appropriate factor. Ox Zn f Zn2+ +2e- Red 2e-+ 4 H+ + 2 VO2+f2 VO2+ + 2 H2O Step 5: Add balanced half-reactions Zn + 4 H+ + 2 VO2+f Zn2+ + 2 VO2+ + 2 H2O

  19. Tips on Balancing Equations • Never add O2, O atoms, or O2- to balance oxygen. • Never add H2 or H atoms to balance hydrogen. • Be sure to write the correct charges on all the ions. • Check your work at the end to make sure mass and charge are balanced. • PRACTICE!

  20. Electrons are transferred from Zn to Cu2+, but there is no useful electric current. CHEMICAL CHANGE fELECTRIC CURRENT With time, Cu plates out onto Zn metal strip, and Zn strip “disappears.” Oxidation: Zn(s) f Zn2+(aq) + 2e- Reduction: Cu2+(aq) + 2e- f Cu(s) -------------------------------------------------------- Cu2+(aq) + Zn(s) f Zn2+(aq) + Cu(s)

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

  22. Zn f Zn2+ + 2e- Cu2+ + 2e- f Cu Oxidation Anode Negative Reduction Cathode Positive • Electrons travel thru external wire. • Salt bridgeallows anions and cations to move between electrode compartments. rAnions Cationsf

  23. The Cu|Cu2+ and Ag|Ag+ Cell

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

  25. Terms Used for Voltaic Cells See Figure 20.6

  26. The Voltaic Pile Drawing done by Volta to show the arrangement of silver and zinc disks to generate an electric current. What voltage does a cell generate?

  27. 1.10 V 1.0 M 1.0 M CELL POTENTIAL, E • Electrons are “driven” from anode to cathode by an electromotive force or emf. • For Zn/Cu cell, this is indicated by a voltage of 1.10 V at 25 ˚C and when [Zn2+] and [Cu2+] = 1.0 M. Zn and Zn2+, anode Cu and Cu2+, cathode

  28. CELL POTENTIAL, E • For Zn/Cu cell, potential is +1.10 V at 25 ˚C and when [Zn2+] and [Cu2+] = 1.0 M. • This is the STANDARD CELL POTENTIAL, Eo • —a quantitative measure of the tendency of reactants to proceed to products when all are in their standard states at 25 ˚C.

  29. Calculating Cell Voltage • Balanced half-reactions can be added together to get overall, balanced equation. Zn(s) f Zn2+(aq) + 2e- Cu2+(aq) + 2e- f Cu(s) -------------------------------------------- Cu2+(aq) + Zn(s) f Zn2+(aq) + Cu(s) If we know Eo for each half-reaction, we could get Eo for net reaction.

  30. CELL POTENTIALS, Eo Can’t measure 1/2 reaction Eo directly. Therefore, measure it relative to a STANDARD HYDROGEN CELL, SHE. 2 H+(aq, 1 M) + 2e- e H2(g, 1 atm) Eo = 0.0 V

  31. Negative electrode Positive electrode Zn/Zn2+ half-cell hooked to a SHE. Eo for the cell = +0.76 V Supplier of electrons Acceptor of electrons Zn f Zn2+ + 2e- Oxidation Anode 2 H+ + 2e- f H2 Reduction Cathode

  32. Reduction of H+ by Zn See Active Figure 20.13

  33. Overall reaction is reduction of H+ by Zn metal. Zn(s) + 2 H+ (aq) f Zn2+ + H2(g) Eo = +0.76 V Therefore, Eo for Zn f Zn2+ (aq) + 2e- is +0.76 V Zn is a (better) (poorer) reducing agent than H2.

  34. Cu/Cu2+ and H2/H+ Cell Eo = +0.34 V Positive Negative Acceptor of electrons Supplier of electrons Cu2+ + 2e- f Cu Reduction Cathode H2f 2 H+ + 2e- Oxidation Anode

  35. Cu/Cu2+ and H2/H+ Cell Overall reaction is reduction of Cu2+ by H2 gas. Cu2+ (aq) + H2(g) fCu(s) + 2 H+(aq) Measured Eo = +0.34 V Therefore, Eo for Cu2+ + 2e- fCu is +0.34 V

  36. + Zn/Cu Electrochemical Cell Zn(s) f Zn2+(aq) + 2e- Eo = +0.76 V Cu2+(aq) + 2e- f Cu(s) Eo = +0.34 V --------------------------------------------------------------- Cu2+(aq) + Zn(s) f Zn2+(aq) + Cu(s) Eo (calc’d) = +1.10 V Anode, negative, source of electrons Cathode, positive, sink for electrons

  37. Uses of Eo Values Organize half-reactions by relative ability to act as oxidizing agents Cu2+(aq) + 2e- f Cu(s) Eo = +0.34 V Zn2+(aq) + 2e- f Zn(s) Eo = –0.76 V Note that when a reaction is reversed the sign of E˚ is reversed!

  38. Uses of Eo Values • Organize half-reactions by relative ability to act as oxidizing agents • Table 20.1 • Use this to predict direction of redox reactions and cell potentials.

  39. Potential Ladder for Reduction Half-Reactions See Figure 20.14 Best oxidizing agents Best reducing agents

  40. Using Standard Potentials, EoTable 20.1 • Which is the best oxidizing agent: O2, H2O2, or Cl2? _________________ • Which is the best reducing agent: Hg, Al, or Sn? ____________________

  41. oxidizing o ability of ion E (V) 2+ Cu + 2e- Cu +0.34 + 2 H + 2e- H 0.00 2+ Zn + 2e- Zn -0.76 reducing ability of element TABLE OF STANDARD REDUCTION POTENTIALS 2

  42. Standard Redox Potentials, Eo Any substance on the right will reduce any substance higher than it on the left. • Zn can reduce H+ and Cu2+. • H2 can reduce Cu2+ but not Zn2+ • Cu cannot reduce H+ or Zn2+.

  43. Cu(s) | Cu2+(aq) || H+(aq) | H2(g) Cathode Positive Anode Negative Electrons r Cu2+ + 2e- f Cu Or Cu f Cu2+ + 2 e- H2f 2 H+ + 2 e- or 2 H+ + 2e- f H2

  44. Cu(s) | Cu2+(aq) || H+(aq) | H2(g) Cathode Positive Anode Negative Electrons r Cu2+ + 2e- f Cu H2f 2 H+ + 2 e- The sign of the electrode in Table 20.1 is the polarity when hooked to the H+/H2 half-cell.

  45. 2+ Cu + 2e- fCu +0.34 + 2 H + 2e- fH2 0.00 2+ Zn + 2e- fZn -0.76 Standard Redox Potentials, Eo Ox. agent Red. agent Any substance on the right will reduce any substance higher than it on the left. • Northwest-southeast rule: product-favored reactions occur between • reducing agent at southeast corner • oxidizing agent at northwest corner

  46. Using Standard Potentials, EoTable 20.1 • In which direction do the following reactions go? • Cu(s) + 2 Ag+(aq) fCu2+(aq) + 2 Ag(s) • Goes right as written • 2 Fe2+(aq) + Sn2+(aq) f2 Fe3+(aq) + Sn(s) • Goes LEFT opposite to direction written • What is Eonet for the overall reaction?

  47. 2+ Cu + 2e- fCu +0.34 + 2 H + 2e- fH2 0.00 2+ Zn + 2e- fZn -0.76 Standard Redox Potentials, Eo CATHODE ANODE • Northwest-southeast rule: • reducing agent at southeast corner = ANODE • oxidizing agent at northwest corner = CATHODE

  48. Standard Redox Potentials, Eo E˚net = “distance” from “top” half-reaction (cathode) to “bottom” half-reaction (anode) E˚net = E˚cathode - E˚anode Eonet for Cu/Ag+ reaction = +0.46 V

  49. Eo for a Voltaic Cell Cd fCd2+ + 2e- or Cd2+ + 2e- fCd Fe fFe2+ + 2e- or Fe2+ + 2e- fFe All ingredients are present. Which way does reaction proceed?

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