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Electrolysis
Chapter 10.1
Electrolysis
• Electrolysis: electrical energy used to cause a
non-spontaneous chemical reaction to occur.
• Electrolytic cell
• All definitions learned previously for galvanic
cells hold. The only difference: the cell voltage
is supplied, not generated.
Electrolysis of molten NaCl
• Why molten?
• 2NaCl(l)  Cl2(g) + 2Na(l)
• anode:
2Cl(l)  Cl2(g) + 2e- Eooxdn = -1.36V
• cathode:
Na+
(l) + e-  Na(l) Eoredn = -2.71V
• Eocell = _________ minimum amt of energy need to
supply
Competing Reactions in Aqueous
Solution
–
e-  Na
Eoredn
= -2.71V
– 2H2O +2 e-  H2(g) + 2OH- Eoredn = -0.83V
– 2H+ + 2e-  H2
• anode:
2Cl-  Cl2(g) + 2e- Eoox = -1.36V
2H2O  O2(g) + 4H+ + 4e- Eoox = -1.23V
• cathode:
Na+ +
• Takes a lot of energy to melt NaCl (801oC)
• Why not use aqueous solution of NaCl?
Eoredn = 0.0V
• The one that you have to supply the least amount of
energy is the one that occurs
• Expect to produce____(g) at anode but
fooled, get ______(g) instead.
• Overvoltage: Difference between
electrode potential and actual voltage
required to cause electrolysis
• Overall:
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Electrolysis of Water
Quantitative Aspects of Electrolysis
• Electrolysis of pure water slow. Why?
• Instead use 0.1M H2SO4 soln
• cathode:
– 2H+ + 2e-  H2
• Measure electrical current needed to supply energy
to cause reactions in amperes (A)
• 1A = 1C/s or 1C = 1A x 1 sec ( 1C is quantity of
electrical charge passing one point in 1s when
current is 1A)
• 1F = 96,500 C = charge carried by 1 mol of e-’s (F =
Faraday)
Eoredn = 0.0V
• anode:
– 2H2O  O2(g) + 4H+ + 4e- Eoox = -1.23V
• Overall 2H2O  2H2(g) + O2(g)
cathode anode
Electroplating
• Active electrodes: electrodes that take part in
electrolysis.
• Example: electrolytic plating.
• Electroplating uses electrolysis to deposit a
thin layer of one metal on another in order to
improve beauty or resistance to corrosion. e.g.
electroplating nickel on a piece of steel.
Electroplating
Electroplating
When aqueous solutions are electrolyzed using metal
electrodes, an electrode will be oxidized if the
oxidation potential is greater than that of water
Ni(s)  Ni2+(aq) + 2e -
E°ox = 0.28
2H2O(l)  4H+(aq) + O2 (g) + 4e - E°ox = -1.23
anode: Ni(s)  Ni2+(aq) + 2e cathode: Ni2+(aq) + 2e -  Ni(s)
Electroplating creates a
silver lining
• Consider an active Ni electrode and another
metallic electrode placed in an aqueous
solution of NiSO4:
• Anode: Ni(s)  Ni2+(aq) + 2e• Cathode: Ni2+(aq) + 2e-  Ni(s).
• Ni plates on the steel electrode.
• Electroplating is important in protecting
objects from corrosion.
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