1. Topic 9: REDOX – Electrochemical cells
Topic 9.2 & Topic 19.1: Voltaic cells convert chemical energy to electrical energy and electrolytic cells convert electrical energy to chemical energy. Energy conversion between electrical and chemical energy lie at the core of electrochemical cells.

2. Important terms for this section
Standard hydrogen electrode (SHE)
Standard conditions
Standard half-cells
Standard electrode potential
Standard reduction potentials
Nernst equation
Electrolyte
Electrodes
Discharge
Inert
Electrochemical cells
Voltaic (galvanic) cells
Electrolytic cells
Half-cells
Electrode potential
Anode
Cathode
Cell diagram convention
Electromotive force (EMF)
Cell potential/electrode potential
Selective discharge
Brine
Electroplating
Galvanized iron
Sacrificial protection

3. Electrochemical cells
Voltaic cells: also called galvanic cells
Uses chemical reaction to generate electricity
Electrolytic cells
Uses electrical energy to drive a chemical reaction

4. Electrochemical cells: voltaic cells
Voltaic cells generate electricity from spontaneous redox RXNs
Example: zinc + copper(II) oxide
Oxidation: Zn(s)  Zn2+(aq) + 2e-
Reduction: Cu2+(aq) + 2e-  Cu(s)
We can separate the reactions physically into half-cells to generate electrode potentials
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5. Electrochemical cells: voltaic cells
The charge separation is called electrode potential
There is an equilibrium of the ions and metal atoms
Where equilibrium lies (and strength of electrode potential) depends on reactivity of metal
Zinc lies toward the ions
Electrode potential is more negative
Copper lies toward the metal atoms
Electrode potential is more positive

6. Electrochemical cells: voltaic cells
Two half-cells, connected, make a voltaic cell
Connected by an external wire can allow e- to flow
E- will flow from the zinc to the copper
Electrodes: anode and cathode
Anode: Zn(s)  Zn2+(aq) + 2e- AnOx (neg charge)
Cathode: Cu2+(aq) + 2e-  Cu(s) RedCat (pos charge)

7. Electrochemical cells: voltaic cells
Needed:
External electronic circuit connected to metal electrodes
Voltmeter to record voltage generated
Salt bridge (glass tube or paper) with aqueous ions
Movement of ions neutralizes buildup of charge on a side and maintains potential difference
Anions move from cathode to anode – opposes flow of e-
Cations move from anode to cathode
Does not interfere with charge, no voltage generated

8. Electrochemical cells: voltaic cells
Cell diagram convention: a shorthand version of drawing a voltaic cell
Single vertical line = physical boundary such as solid electrode and aqueous sol’n
Double vertical lines = salt bridge
Aqueous solutions of electrode next to salt bridge
Anode on left; cathode on right  e- flow from left to right (not universal)
Spectator ions omitted!
If a half-cell contains 2 ions, separate with a comma
Note: e- always flow in external circuit
From anode  cathode

9. Electrochemical cells: voltaic cells
Different half-cells can vary in voltage
Bigger difference in activity series  larger electrode potential  larger voltage generated

10. Electrochemical cells: voltaic cells - summary
Electrons flow from anode to cathode through external circuit
Anions migrate from cathode to anode through salt bridge
Cations migrate from anode to cathode through salt bridge
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12. Electrochemical cells: voltaic cells
Standard electrode potentials
In order to compare half-cell electrode potentials, we need a point of reference
The electromotive force (EMF) = the potential difference in the voltaic cell
Cell potential/electrode potential = E
This is the voltage that depends on the difference in the two half-cells’ ability to undergo reduction
Cannot measure only 1 half-cell. Need the flow of e- to measure the difference
Need a point of reference!!!
Standard Hydrogen Electrode: SHE
This is the baseline for measuring and comparing electrode potentials

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14. Electrochemical cells
Standard Hydrogen Electrode (SHE)
Platinum metal electrode
Inert and won’t ionize
Will catalyze reaction for proton reduction
On surface, platinized platinum (finely divided platinum to increase s.a. and thus RXN rate)
Conditions: 1.0 mol dm-3, pH=0, 100kPa, 278K
The electrode sets up equilibrium of Hydrogen
This is assigned an electrode potential of 0V
Therefore, connecting any other half-cell will give the actual electrode potential

15. Electrochemical cells
Measuring standard electrode potentials
The electrode potentials depend on:
Ion concentrations
Gas pressures
Purity of substance
Temperature
Thus, standard conditions must be maintained: standard half-cells
All solutions have concentration of 1.0 mol dm-3
All gases at pressure 100 kPa
All substances pure
Temperature is 298K
IF the half-cell does not contain a metal, use platinum as the electrode

16. Electrochemical cells
When SHE is connected to another half-cell, the EMF generated is the standard electrode potential, Eθ
Positive Eθ indicates there is greater tendency to reduce than the hydrogen half-cell
Thus, hydrogen half-cell is anode = oxidized
The copper half-cell is cathode = reduced

17. Electrochemical cells
Usually, most metals have Eθ that is negative
Zinc is less reduced than hydrogen
Zinc = anode
Hydrogen = cathode

18. Electrochemical cells: standard reduction potentials
Since there is an equilibrium between reduction and oxidation, there is an agreed convention to standardize the data
Standard electrode potential (SEP) is always given for the reduction reaction
Oxidized species on the left, reduced species on the right
Called standard reduction potentials, Eθ
See section 24 of the data booklet for SEP of half-cells

19. Electrochemical cells: standard reduction potentials
Notes:
All Eθ refer to the reduction reaction
Eθ do not depend on the total number of e-, so are not scaled up or down depending on the stoichiometric equation
The more positive the Eθ for the half-cell, the more readily it is reduced
E- always flow toward the half-cell with the highest Eθ value
In voltaic cells
More negative electrode = anode and oxidation occurs there
More positive electrode = cathode and reduction occurs there

20. Electrochemical cells: standard reduction potentials
How do we apply the use of the standard electrode potential data?
1. We can calculate the cell potential, Eθcell
2. Determine the spontaneity of a reaction
This includes electrode potential and free energy change (Ecell and ∆G)
3. Compare relative oxidizing and reducing power of half-cells

21. Electrochemical cells: standard reduction potentials
1. We can calculate the cell potential, Eθcell
From Eθ of two half-cells, we can calculate EMF
From direction of e- flow, we can predict the outcome of REDOX RXNs
Remember: higher Eθ is reduced (cathode), lower Eθ is oxidized (anode)
Thus,
Notes:
A. Use the reduction potentials for the calculations
B. Even if you multiply the equation, DO NOT multiply the reduction potentials

22. Electrochemical cells: standard reduction potentials
1. We can calculate the cell potential, Eθcell
Example:

23. Electrochemical cells: standard reduction potentials
1. We can calculate the cell potential, Eθcell
Solution:

24. Electrochemical cells: standard reduction potentials
2. Determine the spontaneity of a reaction
This includes electrode potential and free energy change (Ecell and ∆G)
When Eθcell is positive, the reaction is spontaneous
When Eθcell is negative, the reverse reaction is spontaneous

25. Electrochemical cells: standard reduction potentials
2. Determine the spontaneity of a reaction
This includes electrode potential and free energy change (Ecell and ∆G)
n is the number of e- (moles of e-) transferred in the reaction
F is the charge carried by 1 mol e-, also the Faraday constant = 96,500 C mol-1
∆G will be calculated in J so need to convert to kJ

26. Electrochemical cells: standard reduction potentials
3. Compare relative oxidizing and reducing power of half-cells
Metals with low Eθ oxidized easily
More reactive, more likely to lose e-, thus stronger reducing agent
Metal with low Eθ will reduce metal with higher Eθ
Non-metal with high Eθ is reduced easily
More reactive, more likely to gain e-, thus stronger oxidizing agent
Non-metal with high Eθ will oxidize ions non-metal with lower Eθ
In general:
The higher the Eθ, the stronger the oxidizing agent
The lower the Eθ, the stronger the reducing agent

27. Electrochemical cells: standard reduction potentials
3. Compare relative oxidizing and reducing power of half-cells
Example:

28. Electrochemical cells: standard reduction potentials
3. Compare relative oxidizing and reducing power of half-cells

29. Electrochemical cells: Electrolytic cells
Voltaic cells are spontaneous that use chemical E to create electrical E
Electrolytic cells are non-spontaneous and must use an external source to push the e- through to create a chemical breakdown of a reaction
Electrolyte: molten ionic compound or solution of ionic compound
Redox occurs at electrodes when current is passed through electrolyte
Discharge: removal of charges of ions to create neutral products

30. Electrochemical cells: Electrolytic cells

31. Electrochemical cells: Electrolytic cells
Oxidation ALWAYS occurs at the anode
Reduction ALWAYS occurs at the cathode
That being said, the charges are reversed for electrolytic cells

32. Electrochemical cells: determining products in ECs

33. Electrochemical cells: determining products in ECs

34. Electrolytic cells: determining products – molten salts
Molten ionic salts: have very high melting points, so often used in industry
Ex: NaCl to extract Na metal (Aluminum oxide is a common industrial use too)

35. Electrolytic cells: determining products – solutions
Aqueous ionic solutions: have to contend with water ions as well
Predicting the products is a little trickier
Factors to help determine what products will be made:
1. the relative Eθ values for each ion
2. the concentrations of the ions in the electrolyte
3. the type of electrode

36. Electrolytic cells: determining products – water
Electrolysis of water (includes NaOH to help with conductivity)
1. identify ions:
2. write all half-equations

37. Electrolytic cells: determining products – water
Electrolysis of water (includes NaOH to help with conductivity)
3. write balanced equation:
Notice the Na from NaOH is not discharged nor does it interfere with the reaction
4. discuss observed changes:

38. Electrolytic cells: determining products – brine
Electrolysis of brine (NaCl in solution)
1. identify ions:
2. write all half equations:

39. Electrolytic cells: determining products – brine
Electrolysis of brine (NaCl in solution)
3. write balanced equation:
4. observations of reaction:

40. Electrolytic cells: determining products – CuSO4(aq)
Electrolysis of CuSO4(aq)
1. identify ions:
Here, we have 2 options. Use inert electrodes or copper electrodes
Inert: graphite or something else that does not interfere with the reaction
Copper: you can imagine there will be copper everywhere! FUN!

41. Electrolytic cells: determining products – CuSO4(aq)
Electrolysis of CuSO4(aq) – inert electrodes
3. Balanced equation:
4. Observations:

42. Electrolytic cells: determining products – CuSO4(aq)
Electrolysis of CuSO4(aq) – copper electrodes
2. identify ions:
3. balanced equation: net movement of Cu2+(aq) to form Cu(s)
4. observations:
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43. Electrolytic cells: determining products – factors
Q = charge in coulombs, C
I = current in amperes, A
t = time in seconds, s
F = charge carried by 1 mol e- is a Faraday, coulombs

44. Electrolytic cells: determining products – factors
Charge on the ion makes a difference as well
Twice as much electricity is needed to produce 1 mole of Pb than Na

45. Electrolytic cells: electroplating
Electroplating: Coating metal onto another object (metal or conductive)
Features:
1. electrolyte with metal ions wanted for the coating
2. cathode = object to be plated
3. anode = same metal as
coating to replenish electrolyte
Longer time = thicker coating

46. Electrolytic cells: electroplating
Purposes of electroplating
Decorative, it’s pretty – coating other metals, like jewelry, objects, etc.
Corrosion control (aluminum is resistant to corrosion vs. iron)
Sacrificial protection – put zinc on top of iron and it will corrode first
Improve function – chrome on car parts move smoother

47. Electrochemical cells: Summary

48. Electrochemical cells: Summary