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Electrochemistry

Electrochemistry. Chapter 20. IB Topics 9.3-9.5, 19.1 and 19.2 Text Pages 851-882. Electrochemistry. Study of interchange of chemical and electrical energy. Two main processes Generation of electrical current from spontaneous chemical reactions Use of a current to produce chemical change .

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Electrochemistry

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  1. Electrochemistry Chapter 20

  2. IB Topics 9.3-9.5, 19.1 and 19.2 • Text Pages 851-882

  3. Electrochemistry • Study of interchange of chemical and electrical energy. • Two main processes • Generation of electrical current from spontaneous chemical reactions • Use of a current to produce chemical change

  4. Redox Reaction • 8H1+(aq) + MnO41-(aq) +5Fe2+ (aq) Mn2+ (aq) 5Fe3+ (aq) + 4H2O (l) Reduced OxidizedOxidizing ReducingAgentAgent • When the oxidizing and reducing agents are present in the same solution, no useful work is obtained. • Energy is released as heat

  5. Separating the Oxidizing and Reducing Agents of a Redox Reaction Gaining Electrons Losing Electrons Negative charge builds up Positive charge builds up Current flows for an instant then ceases. Charge builds up in the two compartments

  6. Solutions must be connected so that ions can flow to keep net charge in each compartment at zero

  7. Galvanic (Voltaic ) Cell • Device in which chemical energy is converted to electrical energy. • Anode—compartment in which oxidation occurs. • Cathode—compartment in which reduction occurs. • Electrons flow from anode to cathode

  8. Cell Potential/Electromotive force • The reaction in a voltaic cell is always redox. • Oxidizing agent “pulls” electrons through a wire from reducing agent. • Ecell • Emf • Units are volts (V) • Measured by voltmeter

  9. Galvanic (Voltaic) Cells

  10. Practice • 2H1+ (aq) + Zn (s)  Zn2+ (aq)+ H2(g) Al3+ (aq) + Mg (s)  Al (s) + Mg2+ (aq) Show the following for each: -Half reaction at each electrode -Direction of electron flow

  11. Standard Reduction Potentials • The reaction in a voltaic cell is always redox. • 2H1+ (aq) + Zn (s)  Zn2+ (aq)+ H2(g) • Half reactions • Cathode reaction (reduction): 2H1+ (aq) + 2e-  H2(g) • Anode reaction (oxidation) : Zn (s)  Zn2+ (aq) + 2e- Pt electrode is inert. It is in contact with H+(aq). This is called a standard hydrogen electrode.

  12. Standard Reduction Potentials • Total potential of the cell is 0.76 V. • Standard conditions [H+] = 1 M and PH2 = 1atm • E°cell = E°Zn  Zn2+ + E°2H1+  H2 • Red: 2H1+ (aq) + 2e-  H2(g) E°cell = 0 (standard) • Ox: Zn (s)  Zn2+ (aq) + 2e- E°cell = 0.76 V

  13. Standard Reduction Potentials • Half Reactions • Anode: Zn  Zn2+ + 2e- • Cathode: Cu2+ + 2e-Cu • We know that : E° Zn  Zn2+ = 0.76 V • E°cell = E°Zn  Zn2+ + E°Cu2+  Cu • 1.10V = 0.76 V + E°Cu2+  Cu • E°Cu2+  Cu = 0.34 V

  14. Standard Reduction Potentials • Given for reduction half reactions • All solutions are 1 M • All gases are at 1 atm • Using SRP’s for Redox requires the following: • E°cell = E° red cathode- E° red anode • E°cell = E° red cathode+ E° ox anode

  15. Sample Problem 1 • Consider a galvanic (voltaic) cell based on the following reaction • Write the half reactions , assign SRP’s • Give the balanced cell reaction and calculate E°cell Al3+ (aq) + Mg (s)  Al (s) + Mg2+ (aq)

  16. Line Notation for Electrochemical Cells • Shorthand showing two half-cells connected by a salt bridge or porous barrier, : Zn(s)/Zn2+(aq)//Cu2+(aq)/Cu(s) anodecathode The electrodes are shown on the ends and the electrolytes for each side are shown in the middle. 16

  17. Calculating Cell Potentials From Standard Reduction Potenials From the standard reduction table Zn+2 + 2 e-  Zn - 0.763 v Ag+ + e- Ag + 0.799v • Calculate the cell potential for a cell made from silver and zinc electrodes. Since there must be one oxidation and one reduction, the direction of one of two half reactions above must be reversed. Reversing the zinc half reaction making it the oxidation would yield a positive cell potential Zn  Zn+2 + 2 e- +0.763 v (anode) Ag+ + e- Ag + 0.799v (cathode) Cell potential = 1.562 volts 18

  18. Sample Problem 2 Mg2+ + 2e-  Mg E°= -2.37 V Zn2+ + 2e-  Zn E°= - 0.76 V • Does Zn react with Mg2+? • Does Mg react with Zn2+? If Mg is oxidized Mg  Mg2+ + 2e- Eo = +2.37 v Combining this with the reduction of Zn2+ Eo = - 0.76 v Leaves an overall positive cell potential + 1.66 v Therefore Mg reacts with Zn2+. Zn does not react with Mg2+ 19

  19. Metal Displacement Reactions • The electrochemical cell potentials form the basis for predicting which metals will react with salt solution of other metals • This order of reactivity of metals in single replacement reactions is called the activity series • The solid of more reactive metals will displace ions of a less reactive metal from solution. • The relative reactivity of metals is based on potentials of half reactions. • Elements with very different potentials react most vigorously. 20

  20. The Activity Series • Elements with highly negative reduction potentials are not easily reduced but they are easily oxidized. • Since metals react by being oxidized the more negative the reduction potential the more reactive the element. • Elements higher in the table (more negative potential) can displace any element lower (more positive potential). • So Zn + CuCl2 ZnCl2 + Cu Cu + ZnCl2 No Reaction The activity Series is really a reduction potential table arranged from negative to positive 21

  21. Electrolysis • Forcing electricity through a cell to produce a chemical change for which the cell potential is negative. • Causes a non-spontaneous reaction to occur. • Charging batteries, production of aluminum, Chrome plating, obtain reactive metals (Na) from ores

  22. Power source • Anode/Cathode reversed • Electron flow reversed • Ion flow reversed

  23. Electrolytic Cell • 1. Electrons are "produced" in the battery at the anode, the site of oxidation. • 2. The electrons leave the electrochemical cell through the external circuit. • 3. These negative electrons create a negative electrode in the electrolytic cell which attracts the positive Na+ ions in the electrolyte. Na+ ions combine with the free electrons and become reduced (2Na+ + 2e- → Na ) • 4. Meanwhile the negative Cl- become attracted to the positive electrode of the electrolytic cell. At this electrode chlorine is oxidized, releasing electrons (Cl-→ Cl2 + 2 e-) • 5. These electrons travel through the external circuit, returning to the electrochemical cell.

  24. Factors Affecting the Discharge of Ions During Electrolysis • Cation discharged at cathode (negative electrode) • Anion discharged at anode (positive electrode). • In (aq) solutions, H+ and OH- will be present also. • Which ion discharged influenced by • Position in electrochemical series • Concentration • Nature of electrode

  25. Position in Electrochemical Series • The lower the metal ion in the series, the more readily it reduces • Ex: NaOH (aq)— • Hydrogen reduced at negative electrode (cathode) • Sodium oxidized at positive electrode (anode) • Ex: CuSO4— • Copper reduced at negative electrode (cathode) • Hydrogen oxidized at positive electrode (anode)

  26. Concentration • The more concentrated an ion in solution, the greater the production of the substance. • NaCl solution. • Both O2 and Cl2 evolved at positive electrode. • Dilute NaCl (aq) • Mainly O2 evolved. • Concentrated NaCl (aq • Mainly Cl2 evolved.

  27. Nature of the Electrode • Normally safe to assume electrode is inert. • If copper electrodes are used in a CuSO4 solution, the positive electrode will be oxidized to release electons and form Cu2+ while copper is simultaneously reduced at the negative electrode to form copper. The concentration of Cu2+ will remain constant.

  28. Factors Affecting the Quantity of Products • # of electrons • Depends on current and the time it flows. • Time held constant • 2 x current=2 x electrons flowing = 2 X product forming • Current held constant • 2 x time=2 x electrons flowing = 2 X product forming • Charge on ion • Na+(l) +1e-  Na (l) • Formation of one mole of Na requires 1 mole of electrons • Pb2+(l) + 2e-  Pb (l) • Formation of one mole of lead requires 2 moles of elections.

  29. Electroplating • Used in industry to coat one metal with a thin layer of another metal. • Video

  30. Batteries -- Applications of Electrochemical Cells Batteries • device that converts chemical energy into electricity Primary Cells • non-reversible electrochemical cell • non-rechargeable cell Secondary Cells • reversible electrochemical cell • rechargeable cell 31

  31. A Common Dry Cell 32

  32. A 9 Volt Dry Cell 33

  33. “Flash Light” Batteries Dry Cell Zn(s) + 2 MnO2 (s)+ 2 NH4+ (aq)  Zn+2 (aq) + 2 MnO(OH)(s)+ 2 NH3 Alkaline Cell Zn (s) ) + 2 MnO2 (s)ZnO (s)+ Mn2O3(s) 34

  34. Lead-Acid (Car Battery) 35

  35. Lead-Acid (Car Battery) Overall reaction Pb (s) + PbO2 (s) + 2 H2SO4 (aq) = 2 PbSO4 (s) + 2 H2O (l) E = 2.0- volts per cell Cathode PbO2(s) + SO42- (aq) + 4H+ (aq) +2e- PbSO4 (s) + 2 H2O (l) Anode Pb (s) + SO42 (aq) PbSO4 (s) +2e- 36

  36. Nickel-Cadmium (Ni-Cad) Overall reaction Cd(s) + 2 Ni(OH)3(s) = Cd(OH)2(s) + 2 Ni(OH)2(s) E NiCad = 1.25 v/cell Cathode NiO2(s)+ 2 H2O (l) +2e- Ni(OH)2 (s) + 2OH- (aq) Anode Cd(s) + 2OH- (aq) Cd(OH)2 (s) +2e- 37

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