1. Chapter 20 Electrochemistry
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2. Oxidation-reduction reactions Oxidation numbers
Oxidation of metals by acids and salts
The activity series
ALL these to be done in class
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3. Oxidation-Reduction Reactions
In the reaction
Zn(s) + 2H+(aq)  Zn2+(aq) + H2(g).
Which element is oxidised and which one is reduced?
Oxidation – loss of e-
Reduction – gain of e-
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4. Balancing Oxidation-Reduction Reactions
Law of conservation of mass: the amount of each element present at the beginning of the reaction must be present at the end.
Conservation of charge: electrons are not lost in a chemical reaction.
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5. Half Reactions
Half-reactions are a convenient way of separating oxidation and reduction reactions.
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6. Sn2+(aq) + 2Fe3+(aq)  Sn4+(aq) + 2Fe3+(aq)
The half-reactions for
Sn2+(aq) + 2Fe3+(aq)  Sn4+(aq) + 2Fe3+(aq)
are………
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7. Balancing Equations by the Method of Half Reactions
The two incomplete half reactions are
MnO4-(aq)  Mn2+(aq)
C2O42-(aq)  2CO2(g)
Balance the overall reaction equation in an acidic solution
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8. Balancing Equations for Reactions Occurring in Basic Solution
We use OH- and H2O rather than H+ and H2O.
The same method as for acidic solution is used, but OH- is added to “neutralize” the H+ used.
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9. Voltaic Cells
If a strip of Zn is placed in a solution of CuSO4, Cu is deposited on the Zn and the Zn dissolves by forming Zn2+.
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10. Zn is spontaneously oxidized to Zn2+ by Cu2+.
The Cu2+ is spontaneously reduced to Cu0 by Zn.
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12. “Rules” of voltaic cells:
1. At the anode electrons are products. (Oxidation)
2. At the cathode electrons are reagents. (Reduction)
3. Electrons can’t swim!
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14. Anions and cations move through a porous barrier or salt bridge.
Cations move into the cathodic compartment to neutralize the excess negatively charged ions
Anions move into the anodic compartment to neutralize the excess Zn2+ ions formed by oxidation
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15. A Molecular View of Electrode Processes
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17. Cell EMF
e- flow from anode to cathode because the cathode has a lower electrical potential energy than the anode.
1 V is the potential difference required to impart 1 J of energy to a charge of one coulomb:
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18. 1 V is the potential difference required to impart 1 J of energy to a charge of one coulomb:
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19. Cell potential: Ecell is the emf of a cell.
Electromotive force (emf) is the force required to push electrons through the external circuit.
Cell potential: Ecell is the emf of a cell.
For 1M solutions at 25 C (standard conditions), the standard emf (standard cell potential) is called Ecell.
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20. Standard Reduction (Half-Cell) Potentials
Standard reduction potentials, Ered are measured relative to the standard hydrogen electrode (SHE).
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22. For the SHE, we assign 2H+(aq, 1M) + 2e-  H2(g, 1 atm) Ered = 0.
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25. Ecell = Ered(cathode) - Ered(anode) 0.76 V = 0 V - Ered(anode).
For Zn:
Ecell = Ered(cathode) - Ered(anode)
0.76 V = 0 V - Ered(anode).
Therefore, Ered(anode) = V.
Standard reduction potentials must be written as reduction reactions:
Zn2+(aq) + 2e-  Zn(s), Ered = V.
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26. 2Zn2+(aq) + 4e-  2Zn(s), Ered = -0.76 V.
Changing the stoichiometric coefficient does not affect Ered.
Therefore,
2Zn2+(aq) + 4e-  2Zn(s), Ered = V.
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27. The larger the difference between Ered values, the larger Ecell.
Reactions with Ered < 0 are spontaneous oxidations relative to the SHE.
The larger the difference between Ered values, the larger Ecell.
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28. Oxidizing and Reducing Agents
The more positive Ered the stronger the oxidizing agent on the left.
The more negative Ered the stronger the reducing agent on the right.
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30. More generally, for any electrochemical process
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31. Example: For the following cell, what is the cell reaction and Eocell?
Al(s)|Al3+(aq)||Fe2+(aq)|Fe(s)
Al3+(aq) + 3e- → Al(s); EoAl = V
Fe2+(aq) + 2e- → Fe(s); EoFe = V
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32. Example: When an aqueous solution of CuSO4 is electrolysed, Cu metal is deposited:
Cu2+(aq) + 2e- → Cu(s)
If a constant current was passed for 5.00 h and 404 mg of Cu metal was deposited, what was the current?
Ans: 6.81 x 10-2 A
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