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9 & 19. ELECTROCHEMISTRY 1

9 & 19. ELECTROCHEMISTRY 1. Electron Transfer Reactions. 1. Electron transfer reactions are redox reactions. 2. Results in the generation of an electric current (electricity) or be caused by imposing an electric current.

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9 & 19. ELECTROCHEMISTRY 1

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  1. 9 & 19. ELECTROCHEMISTRY 1

  2. Electron Transfer Reactions 1. Electron transfer reactions are redox reactions. 2. Results in the generation of an electric current (electricity) or be caused by imposing an electric current. 3. Therefore, this field of chemistry is often called ELECTROCHEMISTRY.

  3. Two industrial applications: • I. Voltaic Cells ____ Produces energy from a spontaneous chemical reaction. • II. Electrolytic Cells____ Uses energy in order to promote a non spontaneous chemical reaction.

  4. Voltaic Cells/ Electrochemical Cells A devicethatobtainselectricalenergyfrom a spontaneouschemicalreaction Batteries are voltaic cells

  5. Terms Used for Voltaic Cells

  6. Zn --> Zn2+ + 2e- Cu2+ + 2e- --> Cu Oxidation Anode Negative Reduction Cathode Positive •Electrons travel thru external wire. • Salt bridge allows anions and cations to move between electrode compartments. <--Anions Cations--> RED CAT

  7. AN OX chases a RED CAT

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

  9. Eo for a Voltaic Cell Cd --> Cd2+ + 2e- or Cd2+ + 2e- --> Cd Fe --> Fe2+ + 2e- or Fe2+ + 2e- --> Fe All ingredients are present. Which way does reaction proceed?

  10. CHEMICAL CHANGE --->ELECTRIC CURRENT With time, Cu plates out onto Zn metal strip, and Zn strip “disappears.” • Zn is oxidized and is the reducing agent Zn(s) ---> Zn2+(aq) + 2e- • Cu2+ is reduced and is the oxidizing agentCu2+(aq) + 2e- ---> Cu(s)

  11. More About Calculating Cell Voltage 2 H2O + 2e- ---> H2 + 2 OH- Cathode 2 I----> I2 + 2e- Anode ------------------------------------------------- 2 I- + 2 H2O --> I2 + 2 OH- + H2 Assume I- ion can reduce water. Assuming reaction occurs as written, E˚ = E˚cat+ E˚an= (-0.828 V) - (- +0.535 V) = -1.363 V Minus E˚ means rxn. occurs in opposite direction (the connection is backwards or you are recharging the battery)

  12. http://ibchem.com/IB/ibnotes/full/red_htm/19.2.htm • Write the equation for the spontaneous reaction that will occur when a magnesium half cell is connected to na aluminum half cell.

  13. 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.

  14. Dry Cell Battery Anode (-) Zn ---> Zn2+ + 2e- Cathode (+) 2 NH4+ + 2e- ---> 2 NH3 + H2

  15. Alkaline Battery Nearly same reactions as in common dry cell, but under basic conditions. Anode (-): Zn + 2 OH- ---> ZnO + H2O + 2e- Cathode (+): 2 MnO2 + H2O + 2e- ---> Mn2O3 + 2 OH-

  16. Mercury Battery Anode: Zn is reducing agent under basic conditions Cathode: HgO + H2O + 2e- ---> Hg + 2 OH-

  17. Lead Storage Battery Anode (-) Eo = +0.36 V Pb + HSO4- ---> PbSO4 + H+ + 2e- Cathode (+) Eo = +1.68 V PbO2 + HSO4- + 3 H+ + 2e- ---> PbSO4 + 2 H2O

  18. Ni-Cad Battery Anode (-) Cd + 2 OH- ---> Cd(OH)2 + 2e- Cathode (+) NiO(OH) + H2O + e- ---> Ni(OH)2 + OH-

  19. H2 as a Fuel Cars can use electricity generated by H2/O2 fuel cells. H2 carried in tanks or generated from hydrocarbons

  20. Balancing Equations for Redox Reactions Some redox reactions have equations that must be balanced by special techniques. MnO4- + 5 Fe2+ + 8 H+ ---> Mn2+ + 5 Fe3+ + 4 H2O Mn = +7 Fe = +2 Mn = +2 Fe = +3

  21. Balancing Equations Consider the reduction of Ag+ ions with copper metal. Cu + Ag+ --give--> Cu2+ + Ag

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

  23. 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 ---> Cu2+ + 2e- Oxidizing agent 2 Ag+ + 2 e- ---> 2 Ag Step 5: Add half-reactions to give the overall equation. Cu + 2 Ag+ ---> Cu2+ + 2Ag The equation is now balanced for both charge and mass.

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

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

  26. 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!

  27. 1. Can an acidified aqueous solution of potassium dichromate spontaneously oxidize a solution of bromide ions to bromine?

  28. 2. Can a solution of tin II ions reduce a solution of iron III ions? If so, are the iron III ions reduced to iron II or to iron metal?

  29. 3. What will happen when copper I sulfate(s) dissolves in water?

  30. II. Electrolysis http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch20/faraday.php Electrolysis is the situation when redox cells are forced to run in reverse by attaching an electricity source to overcome the potential difference.

  31. Salgema-Maceio/Brasil

  32. Some applications:

  33. a)Electrolysis of molten NaCl: NaCl(l) Species present: Na+ and Cl- electrolyte

  34. Electrolysis of molten NaCl: NaCl(l) Species present: Na+ and Cl- Positive electrode=> oxidation Anode 2 Cl- (l) => Cl2 (g) + 2e Negative electrode=> reduction Cathode Na+ (l)+ e => Na(l) Overall: 2 Cl-(l)+ 2 Na+ (l)=>Cl2 (g) + 2 Na(l) electrolyte

  35. b)Electrolysis of aqueous NaCl, NaCl(aq) Species: Na+1 , H+1 , Cl-1 , H2O This time,H+ will be reduced instead of Na+ Cathode (-): Reduction 2 H2O(l) + 2 e- =>H2(g) + 2 OH-(aq) Anode (+): Oxidation   2 Cl- => Cl2 (g)+ 2 e- Overall: 2 NaCl(aq) + 2 H2O(l) => 2 Na+(aq) + 2 OH-(aq) + H2(g) + Cl2(g)

  36. Because the demand for chlorine is much larger than the demand for sodium, electrolysis of aqueous sodium chloride is a more important process commercially. • Electrolysis of an aqueous NaCl solution has two other advantages: It produces H2 gas at the cathode, which can be collected and sold. It also produces NaOH, which can be drained from the bottom of the electrolytic cell and sold.

  37. Electrolysis Al

  38. 9.5.3. How electric current is conducted. • http://www.youtube.com/watch?v=Y9qMR3GV7WA • In an electrolytic cell, current is conducted by electrons in the wire and by ions in the electrolyte.

  39. Hodder Q5

  40. II. Electrolytic Cell The ions that are successfully released at the electrodes depend on three factors 1)The position of the ion in the electrochemical series. As a rule of thumb, if the metal appears below hydrogen in the electrochemical series then it will be preferentially deposited. 2)The concentration of the ion in the solution. 3)The nature of the electrode: Platinum, graphite • http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch20/faraday.php

  41. 9.5.4. Deduce the products of the electrolysis of any molten salt, PbBr2 • http://www.youtube.com/watch?v=kINjUBolU3M&feature=related • Hodder Q1b and Q2

  42. The ions that are successfully released at the electrodes depend on three factors The ions that are successfully released at the electrodes depend on three factors 1)The position of the ion in the electrochemical series. As a rule of thumb, if the metal appears below hydrogen in the electrochemical series then it will be preferentially deposited. 2)The concentration of the ion in the solution. 3)The nature of the electrode: Platinum, graphite • http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch20/faraday.php

  43. IB Questions

  44. http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch20/faraday.phphttp://chemed.chem.purdue.edu/genchem/topicreview/bp/ch20/faraday.php • Because the demand for chlorine is much larger than the demand for sodium, electrolysis of aqueous sodium chloride is a more important process commercially. • Electrolysis of an aqueous NaCl solution has two other advantages. It produces H2 gas at the cathode, which can be collected and sold. It also produces NaOH, which can be drained from the bottom of the electrolytic cell and sold.

  45. Solid sodium chloride doesn't conduct electricity, because there are no electrons which are free to move. When it melts, sodium chloride undergoes electrolysis, which involves conduction of electricity because of the movement and discharge of the ions. In the process, sodium and chlorine are produced. This is a chemical change rather than a physical process.

  46. Predict the products of electrolysis of strong calcium chloride solution • At the cathode • Species present Ca2+ and H+. Ca2+ is higher in the reactivity seriers than hydrogen and therefore cannot be released. • The reaction is therefore: 2H+(aq) + 2e H2(g) • At the anode • Species present OH- and Cl- . The chloride concentration is strong and so it is preferentially oxidised and the reaction is: 2Cl-(aq) Cl2(g) + 2e • Species remaining in solution: Calcium ions and hydroxide ions

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