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Electrochemistry

Electrochemistry. Making chemistry useful!. Oxidation and Reduction. Oxidation – the loss of electrons (increase charge) Reduction – the gain of electrons (decrease of charge). Is Redox Occurring?. Oxidation Numbers – a bookkeeping method to keep track of electrons.

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Electrochemistry

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  1. Electrochemistry Making chemistry useful!

  2. Oxidation and Reduction • Oxidation – the loss of electrons (increase charge) • Reduction – the gain of electrons (decrease of charge)

  3. Is Redox Occurring? • Oxidation Numbers – a bookkeeping method to keep track of electrons. Is this redox? H2O + Al + MnO4- Al(OH)4- + MnO2

  4. Assigning Oxidation Numbers • If it is in its elemental form ON = 0! • Monatomic Ions have ON = charge of ion • Oxygen: ON = -2 unless O22- (peroxide) • Hydrogen: ON = -1 (with metals) ON = +1 (with non metals) • Fluorine: ON = -1 (always) • Sum of Oxidation Numbers must equal the net charge.

  5. Practice Assigning Oxidation Numbers • On a whiteboard, assign oxidation numbers to the following species: H2 Mg As3- Ca2+ MnO4- SCl2 SO42- C2O42- Al(OH)4- MnO2

  6. Is oxidation occurring? • Do the oxidation numbers change? Are these redox? H2O + Al + MnO4- Al(OH)4- + MnO2 Cd + NiO2 + H2O Cd(OH)2 + Ni(OH)2 Pb(NO3)2 + KI  PbI2 + KNO3

  7. The Agents! • The oxidizing agent IS REDUCED! • The reducing agent IS OXIDIZED! • This is because redox cannot occurr separately! If something oxidizes something else (oxidizing agent) it must be reduced! • There is no electron purgatory!

  8. Balancing Redox Reactions • Sometimes it is easy! • 2H2O + Al + MnO4- Al(OH)4- + MnO2 • Sometimes it is not! • MnO4- + C2O42- Mn2+ + CO2 Some hydrogens from water are needed to balance! 16H+ + 2MnO4- + 5C2O42- 2Mn2+ + 8H2O + 10CO2

  9. Balancing Redox Reactions…cont’d • Since water is always present we can use them in the reaction. • So how the heck do we do that? • We create rules! • Acidic Solution • Basic Solution

  10. Balancing in Acidic Solution • Divide into half reactions. • Balance the half reactions: • Balance the NON-hydrogen nor oxygen • Balance the Oxygen with H2O • Balance the H with H+ (since it is acidic) • Balance the charge with electrons (e-) • Equalize the electrons • Add the half reactions

  11. Acidic Balancing…an example! • Step 1: Divide into half reactions MnO4- + C2O42- Mn2+ + CO2 Put all the non-O and non-H that are oxidized together (same for reduced) MnO4- Mn2+ C2O42- CO2 ON = +7 ON = +2 REDUCTION! – ON goes down ON = +3 ON = +4 OXIDATION! – ON goes up

  12. Acidic Balancing…an example! • Step 2: Balance the half reactions! • Balance the Non O and H MnO4- Mn2+ C2O42- 2 CO2 • Balance the O with water MnO4- Mn2+ + 4 H2O C2O42- 2 CO2

  13. Acidic Balancing…an example! • Step 2: Balance the half reactions! • Balance the H with H+ 8 H+ + MnO4- Mn2+ + 4 H2O C2O42- 2 CO2 • Balance charge with e- 5 e- + 8 H+ + MnO4- Mn2+ + 4 H2O C2O42- 2 CO2 + 2 e-

  14. Acidic Balancing…an example! • Step 3: Equalizing the reactions • 2[5 e- + 8 H+ + MnO4- Mn2+ + 4 H2O ] 5[C2O42- 2 CO2 + 2 e-] There are TEN electrons transferred in this reaction!! Make the electrons equal since oxidation and reduction must occur together!

  15. Acidic Balancing…an example! • Step 4: Add the half reactions 10 e-+16 H+ + 2MnO4-2 Mn2+ + 8 H2O 5C2O42-10 CO2 + 10 e- 16H+ + 2MnO4- + 5C2O42-2Mn2+ + 10CO2 + 8H2O Cross off what occurs on BOTH sides of the arrow! Note the water and hydrogen ions do contribute to the reaction since they are present in an acidic solution.

  16. Acidic Balancing…You try it! • Use the steps to balance the following equations on your white board: Cr2O72- + Cl- Cr3+ + Cl2 Mn2+ + BiO3-  Bi3+ + MnO4-

  17. Cr2O72- + Cl- Cr3+ + Cl2Mn2+ + BiO3-  Bi3+ + MnO4- • Divide into half reactions. • Balance the half reactions: • Balance the NON-hydrogen nor oxygen • Balance the Oxygen with H2O • Balance the H with H+ (since it is acidic) • Balance the charge with electrons (e-) • Equalize the electrons • Add the half reactions

  18. Basic Balancing…what’s the difference? • After you add the H+, you must cancel it! There are no H+ in solution since it is basic! • But there is lots of OH-!! So add it to both sides to cancel out the H+ • HINT: What is OH- +H+ ?? • Practice: NO2- + Al  NH3 + Al(OH)4-

  19. Voltaic Cells: the workhorses of electrochemistry! • The engineering of electrochemistry! • Electrons are transferred (called electricity) • Set up the redox reaction to do electrical work!

  20. NOT Useful! • Watch these not so useful redox reactions: Video #1 Video #2 Video #3

  21. Let’s Make it More Useful • Watch the video. • Important Parts of a voltaic cell: • Cathode (Reduction) • Anode (Oxidation) • Circuit (electron path) • Salt Bridge (even out charge) • This is accomplished by SEPARATING the oxidation and reduction.

  22. Voltaic Cell: a picture Salt Bridge Anode Cathode Circuit Since the reactions are separated, the electrons are forced to flow through the circuit to create a reaction. We use them as they go by!

  23. Voltaic Cell: where do things occur? Electrons leave the anode since it is negative. Electrons go to the cathode since it is positive. Cations go to where the electrons do to balance the charge and vice versa for anions.

  24. Calculating the Cell Potential • We know that some elements like to be oxidized more than others:

  25. How much difference in potential? • The difference between the desire to be oxidized between the cathode and the anode will tell us how much potential there is for each voltaic cell: Eocell = Eocathode – Eoanode (reduction – oxidation)

  26. Standard REDUCTION Potentials Likes to be reduced, positive REDUCTION potential • For convention, we write all reactions as REDUCTION, regardless of which they would rather do. Likes to be oxidized, negative REDUCTION potential

  27. Standard to whom? • HYDROGEN! • Why? Because! Hydrogen is the random chosen zero value. All are compared to it!

  28. Who will oxidize? • The one higher on the list will reduce better than hydrogen and the one lower on the list would rather oxidize. • Higher one will be the cathode and the lower one will be the anode. • The difference between them is the cell potential.

  29. Standard Cell Potential • The difference between the desire to be oxidized between the cathode and the anode will tell us how much potential there is for each voltaic cell: Eocell = Eocathode – Eoanode (reduction – oxidation)

  30. Calculate Cell Potential • Zn + Cu2+ Zn2+ + Cu • Split to half reactions: • Zn  Zn2+ (rather oxidize, lower, anode) • Cu2+  Cu (rather reduce, higher, cathode) • EoZn = -0.76 V EoCu = +0.34 V • Eocell = Eocathode – Eoanode = +0.34 – (-0.76) = 1.10 V

  31. You Practice! • P 822 # 24 • P 822 # 22 • PS – what does “standard” cell potential mean?

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