1. Galvanic Cell or Voltaic

2. Electrolytic Cells YOU ARE EXPECTED TO BE ABLE TO:
Construct a labelled diagram to show the structure of an electrolytic and electrochemical cell and describe its operation
Identify half reactions that take place at the anode and cathode of each cell
Predict the most likely products
Calculate the minimum voltage necessary for the electrolysis of a solution under standard conditions
Carry out calculations related to the current passing through an electrolytic cell and the quantities of product formed

3. CHEMICAL CHANGE  ELECTRIC CURRENT
To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs thru an external wire.
This is accomplished in a VOLTAIC cell.
(also called GALVANIC cell) technically these are not batteries.
A group of such cells is called a battery.

4. CELL POTENTIAL, E0
Zn  Zn2+ + 2e-
ANODE
2e- + Cu2+ Cu
CATHODE
Electrons are “driven” from anode to cathode by an electromotive force or emf.
For Zn/Cu cell, a voltage of 1.10 V at 25°C and when [Zn2+] and [Cu2+] = 1.0 M.

5. Electrochemical Cell Electrons move from anode to cathode in the wire.
Anions & cations move through the salt bridge.
Electrochemical Cell

6. Electrochemical Cells
Cell Diagram
Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq)
[Cu2+] = 1 M & [Zn2+] = 1 M
Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)
anode
cathode

7. Electrochemical Cells
spontaneous
redox reaction

8. A voltaic cell based on the zinc-copper reaction
oxidation half-reaction:
Zn(s) Zn2+(aq) + 2e-
reduction half-reaction:
Cu2+(aq) + 2e Cu(s)
overall (cell) reaction:
Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

9. Zn2+(aq)
Zn2+(aq) e–  Zn(s)

12. Zn  Zn2+ + 2 e- oxidation Cu2+ + 2 e-  Cu reduction
- electrode
anode
oxidation
+ electrode
cathode
reduction
electron flow
At this electrode the metal loses electrons and so is oxidised to metal ions.
These electrons make the electrode negative.
As CATIONS (+) migrate through the salt bridge electrode is positive. Zn Cu
Zn  Zn e-
oxidation
Cu e-  Cu
reduction

13. STANDARD REDUCTION POTENTIALS
Oxidizing ability of ion
Half-Reaction Eo (Volts)
Reducing ability
of element
Cu e-  Cu
2 H+ + 2e-  H
Zn e-  Zn
BEST Oxidizing agent ? ?
Cu2+
BEST Reducing agent ? ? Zn

14. Standard Notation for Electrochemical Cells
Phase boundary
Salt bridge
Phase boundary
ANODE Zn / Zn2+ // Cu2+ / Cu CATHODE
Cathode electrode
Anode electrode
Active electrolyte in
reduction half-reaction
Active electrolyte in
oxidation half-reaction
OXIDATION
REDUCTION

16. Notation for a voltaic cell
components of anode compartment
(oxidation half-cell)
components of cathode compartment
(reduction half-cell)
phase of lower oxidation state
phase of higher oxidation state
phase of higher oxidation state
phase of lower oxidation state
phase boundary between half-cells
examples:
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu (s)
Zn(s) Zn2+(aq) + 2e-
Cu2+(aq) + 2e Cu(s)
graphite | I-(aq) | I2(s) || H+(aq), MnO4-(aq) | Mn2+(aq) | graphite
inert electrode

17. ½ Equations and e- flow University Berkley (6 min)

18. Standard Conditions Concentration
1.0 mol dm-3 (ions involved in ½ equation)
Temperature
298 K
Pressure
100 kPa (if gases involved in ½ equation)
Current
Zero (use high resistance voltmeter)

19. S tandard
H ydrogen
E lectrode

20. University of Berkley S.H.E. 3 (min)

22. Pt(s) | H2(g) | H+(aq) || Cu2+(aq) | Cu(s)

25. The ionic compound is called an electrolyte.
We have learnt that ionic compounds conduct electricity when molten or aqueous.
When an electric current passes through such compounds, the compounds are decomposed in a chemical reaction. This is known as electrolysis.
The ionic compound is called an electrolyte.
In this lesson, we will learn about the electrolysis of
Molten ionic compounds
Aqueous ionic compounds 4

26. Electrolytic Cells
In an electrolytic cell, a non-spontaneous redox reaction is made to occur by pumping electrical energy into the system DC
CATHODE -
+ ANODE
Reduction occurs at the cathode
M+ + e-  M
Oxidation occurs at the anode
X-  X + e-

27. Voltaic v Electrolytic cell
voltaic cell
electrolytic cell
oxidation half-reaction:
Sn(s) Sn2+(aq) + 2e-
oxidation half-reaction:
Cu(s) Cu2+(aq) + 2e-
reduction half-reaction:
Cu2+(aq) + 2e Cu(s)
reduction half-reaction:
Sn2+(aq) + 2e Sn(s)
overall (cell) reaction
Sn(s) + Cu2+(aq) Sn2+(aq) + Cu(s)
overall (cell) reaction:
Sn(s) + Cu2+(aq) Sn2+(aq) + Cu(s)

28. Voltaic V Electrolytic cells
VOLTAIC CELL
ELECTROLYTIC CELL
Energy is released from spontaneous redox reaction
system does work on its surroundings
Energy is absorbed to drive a non-spontaneous redox reaction
surroundings (power supply)
do work on the system (cell)
oxidation half-reaction:
X X+ + e-
oxidation half-reaction:
A A + e-
reduction half-reaction:
reduction half-reaction:
B+ + e B
Y+ + e Y
overall (cell) reaction
X + Y X+ + Y
overall (cell) reaction:
A- + B A + B

29. Electrolysis describes what happens in an electrolytic cell and
means to use electricity to make chemicals.
Many elements are made by electrolysis
Pb Al Zn Na K Li H2 Cl2 F2 I2 O2
Pb e- → Pb(s)
2Cl → Cl2(g) + 2e-
Sometimes called electrowinning- the element is won from its ion

30. Applying Faraday’s Laws
1 mole e- - +
AgNO3 solution
cathode
CuSO4 solution
cathode
1 mole Ag (108g) deposited on cathode
½ mole Cu (32g) deposited on cathode

31. Making Aluminum by Electrolysis
Alcan (70,000 employees in 55 countries)
Kitimat B.C MT 7 % world production

32. Aluminum Production by Electrolysis
Name of the Ore imported from (Guinea and Brazil)
Bauxite
Al2O3.3H2O
Heating drives off the water
Al2O3.3H2O + Heat → Al2O3 + 3H2O
Melting point of Bauxite is C
This is too hot!
Cryolite is added
Lowers the melting point to C

33. Reduction of water
You cannot reduce
Aluminum in water!
It must be molten!

34. Liquid Al floats to the top and is removed
DC Power
Oxygen gas C C
Reduction
Cathode is -
Al3+ + 3e- → Al(s)
Oxidation
Anode +
O2- → 1/2O2(g) + 2e-
Al3+
O2-
Al2O3(l)
Cation Cathode Reduction
Anion Anode Oxidation

35. Electroplating e- Reduction Cathode is - Au Au+ + e- → Au(s) Oxidation
Au plating a Cu penny
Electrolyte: Must contain the ion of the metal that plates
Cathode: The metal to be covered with a new metal
Anode: Metal to be plated on top the other metal
DC Power e-
Reduction
Cathode is -
Au+ + e- → Au(s) Au
Oxidation
Anode is +
Au(s) → Au e-
Au+
CN- Cu Cu

36. +ve
-ve
stainless steel or Au
AuCN

37. Au plated

38. Copper Ring
Gold Plated

39. Electroplating e- Reduction Ag Cathode - Ag+ + e- → Ag(s) Oxidation
Ag plating a Loonie
DC Power e-
Reduction
Cathode -
Ag+ + e- → Ag(s) Ag
Oxidation
Anode +
Ag(s) → Ag e-
Ag+
NO3- $1 $1

40. Stoichiometry with Faraday’s Constant
A technician is plating a faucet with 0.86 g of Cr from an electrolytic bath containing aqueous Cr2(SO4)3. If 12.5 min is allowed for the plating, what current is needed?
SOLUTION:
Cr3+(aq) + 3e Cr(s)
0.86 g (mol Cr) (3 mol e-)
(52.00 g Cr) (mol Cr)
= mol e-
0.050 mol e- (9.65 x 104 C/mol e-) = 4.8 x 103 C
4.8 x 103 C
12.5 min
(min)
(60 s) x
= 6.4 C/s = 6.4 A x

41. Batteries Dry cell Leclanché cell Anode: Zn (s) Zn2+ (aq) + 2e-
Cathode:
2NH4 (aq) + 2MnO2 (s) + 2e Mn2O3 (s) + 2NH3 (aq) + H2O (l) +
Zn (s) + 2NH4 (aq) + 2MnO2 (s) Zn2+ (aq) + 2NH3 (aq) + H2O (l) + Mn2O3 (s)

42. Non-rechargeable (primary) cells – Zinc-carbon
-0.80 V Zn(NH3) e-  Zn NH3
+0.70 V 2 MnO H e-  Mn2O3 + H2O
Standard cell
Short life
Determine: a) cell emf
b) overall reaction during discharge

43. Non-rechargeable (primary) cells – alkaline
-0.76 V Zn e-  Zn
+0.84 V MnO2 + H2O + e-  MnO(OH) + OH-
Determine: a) cell emf
b) overall reaction during discharge
Longer life

44. Batteries Mercury Battery Anode:
Zn(Hg) + 2OH- (aq) ZnO (s) + H2O (l) + 2e-
Cathode:
HgO (s) + H2O (l) + 2e Hg (l) + 2OH- (aq)
Zn(Hg) + HgO (s) ZnO (s) + Hg (l)

45. Rechargeable (secondary) cells
In non-rechargeable (primary) cells, the chemicals are used up so the voltage drops
In rechargeable (secondary) cells the reactions are reversible – they are reversed by applying an external current.
It is important that the products from the forward reaction stick to the electrodes and are not dispersed into the electrolyte.

46. Rechargeable (secondary) cells – Li ion
+0.60 V Li+ + CoO2 + e-  LiCoO2
-3.00 V Li+ + e-  Li
Rechargeable
Most common rechargeable cell
Determine: a) cell emf
b) overall reaction during discharge
c) overall reaction during re-charge

47. Rechargeable (secondary) cells – lead-acid
+1.68 V PbO2 + 3 H+ + HSO e-  PbSO4 + 2 H2O
-0.36 V PbSO4 + H e-  Pb + HSO4-
Determine: a) cell emf
b) overall reaction during discharge
c) overall reaction during re-charge
Used in sealed car batteries (6 cells giving about 12 V overall)

48. Rechargeable (secondary) cells – nickel-cadmium
+0.52 V NiO(OH) + 2 H2O + 2 e-  Ni(OH)2 + 2 OH-
-0.88 V Cd(OH) e-  Cd OH-
Determine: a) cell emf
b) overall reaction during discharge
c) overall reaction during re-charge

49. Batteries
A FUEL CELLis an electrochemical cell that requires a continuous supply of reactants to keep functioning
Anode:
2H2 (g) + 4OH- (aq) H2O (l) + 4e-
Cathode:
O2 (g) + 2H2O (l) + 4e OH- (aq)
2H2 (g) + O2 (g) H2O (l)

50. Pros & cons of cells
+ portable source of electricity
Pros & cons of non-rechargeable cells
+ cheap, small
– waste issues
Pros & cons of rechargeable cells
+ less waste, cheaper in long run
– still some waste issues
Pros & cons of fuel cells
+ water is only product
– most H2 is made using fossil fuels, fuels cells expensive