1. Primary Cells
A primary cell is a cell that can be used once only and cannot be recharged.
The reactants cannot be regenerated

2. Primary cells non-rechargeable These cells are not rechargeable.
․Zinc-carbon cells
Recharging is dangerous as it produces H2 and heat which results in an explosion.

3. Primary cells These cells are not rechargeable.
․Alkaline manganese cells

4. Primary cells These cells are not rechargeable.
․Silver oxide cells (button cells)

5. Primary cells These cells are not rechargeable.
․Lithium primary cells (button cells)

6. Zinc-carbon Cells

7. Zinc-carbon Cells At anode: Zn(s) Zn2+(aq) + 2e– At cathode:
2MnO2(s) + 2NH4+(aq) + 2e– Mn2O3(s) + 2NH3(aq) + H2O(l)
Overall reaction:
Zn(s) + 2MnO2(s) + 2NH4+(aq)  Zn2+(aq) + Mn2O3(s) + 2NH3(aq) + H2O(l)
Ecell = V

8. Zinc-carbon Cells The cell diagram for the zinc-carbon cell is:
Zn(s) | Zn2+(aq) [2MnO2(s) + 2NH4+(aq)], [Mn2O3(s) + 2NH3(aq) + H2O(l)] | C(graphite)
Overall reaction:
Zn(s) + 2MnO2(s) + 2NH4+(aq)  Zn2+(aq) + Mn2O3(s) + 2NH3(aq) + H2O(l)

9. At anode,
Zn(s) + 2OH(aq) ZnO(s) + H2O(l) + 2e
At cathode,
Ag2O(s) + H2O(l) + 2e Ag(s) + 2OH(aq)
Overall reaction :Zn(s) + Ag2O(s)  ZnO(s) + 2Ag(s)

10. Q.20(a) At anode, Zn(s) + 2OH(aq) ZnO(s) + H2O(l) + 2e At cathode,
HgO(s)
At anode,
Zn(s) + 2OH(aq) ZnO(s) + H2O(l) + 2e
At cathode,
HgO(s) + H2O(l) + 2e Hg(l) + 2OH(aq)
Overall reaction : Zn(s) + HgO(s)  ZnO(s) + Hg(l)

11. Q.20(b)
HgO(s)
= V – (1.216V) = 1.314V

12. Secondary Cells Electrochemical cells that can be recharged.
Examples : -
Lead-acid accumulators
Nickel-cadmium cells (NiCad)
Nickel-Metal hydride(NiMH) cells
Lithium-ion cells

13. Lead grids coated with PbSO4(s)
Pb(s) + H2SO4(aq)  PbSO4(s) + H2(g)

14. spongy lead Lead grids coated with PbSO4(s) During charging
Negative electrode :
PbSO4(s) + 2e Pb(s) + SO42(aq)
spongy lead

15. spongy PbO2 Lead grids coated with PbSO4(s) During charging
Positive electrode :
PbSO4(s) + 2H2O(l)
PbO2(s) + 4H+(aq) + SO42(aq) + 2e
spongy PbO2

16. Lead grids coated with PbSO4(s)
During discharging
Anode :
Pb(s) + SO42(aq) PbSO4(s) + 2e

17. Lead grids coated with PbSO4(s)
During discharging
Cathode :
PbO2(s) + 4H+(aq) + SO42(aq) + 2e PbSO4(s) + 2H2O(l)

18. Overall reaction : -
Pb(s) + PbO2(s) + 2H2SO4(aq) PbSO4(s) + 2H2O(l)
discharge
charge
The cell diagram for the lead-acid accumulator is:
Pb(s) | PbSO4(s) [PbO2(s) + 4H+(aq) + SO42–(aq)], [2PbSO4(s) + 2H2O(l)] | Pb(s)

19. Overall reaction : - PbSO4 is coated on the electrodes,
Pb(s) + PbO2(s) + 2H2SO4(aq) PbSO4(s) + 2H2O(l)
discharge
charge
PbSO4 is coated on the electrodes,
The reversed processes are made possible.

20. Overall reaction : -
Pb(s) + PbO2(s) + 2H2SO4(aq) PbSO4(s) + 2H2O(l)
discharge
charge
The cell should be charged soon after complete discharge
Otherwise, fine ppt of PbSO4 will become coarser and inactive, making the reversed process less efficient.

21. Overall reaction : - Pb(s) and PbO(s) are on different electrodes
Pb(s) + PbO2(s) + 2H2SO4(aq) PbSO4(s) + 2H2O(l)
discharge
charge
charge
Pb(s) and PbO(s) are on different electrodes
Direct reaction is not possible
Porous partition is not needed

22. Overall reaction : - During discharging, H2SO4 is being used up
Pb(s) + PbO2(s) + 2H2SO4(aq) PbSO4(s) + 2H2O(l)
discharge
charge
During discharging, H2SO4 is being used up
The density of electrolyte solution 
The charging/discharging status can be monitored by a hydrometer.

23. Q.21
Ecell = Eocathode – Eoanode
= (1.69V) – (0.35V) = 2.04V

24. Nickel-cadmium cells – Nicad cells
Q.22(a)
At anode,
Cd(s) + 2OH(aq) Cd(OH)2(s) + 2e
At cathode,
NiO(OH)(s) + H2O(l) + e
Ni(OH)2(s) + OH(aq)

25. Nickel-cadmium cells – Nicad cells
Q.22(b)
Overall reaction : -
2NiO(OH)(s) + Cd(s) + 2H2O(l) 
2Ni(OH)2(s) + Cd(OH)2(s)

26. Nickel metal hydride cell
(NiMH)
Cathode : NiO(OH)
Anode : MH(s)
where M is a hydrogen-absorbing alloy.
More environmentally friendly than NiCad cell due to the absence of Cd.

27. Nickel metal hydride cell (NiMH)
On discharging,
Anode : -
MH(s) + OH(aq) M(s) + H2O(l) + e +1
At cathode,
NiO(OH)(s) + H2O(l) + e
Ni(OH)2(s) + OH(aq) +3 +2

28. Nickel metal hydride cell
(NiMH)
Voltage : 1.2 V
Electrolyte : KOH
2 to 3 times the capacity of an equivalent NiCad cell
From 1100 mAh up to 8000 mAh.

29. Anode is graphite into which Li+ are inserted
Lithium ion cell

30. Cathode is metal oxide into which Li+ are inserted.
Lithium ion cell

31. During charging, Li+ moves from cathode to anode
Lithium ion cell

33. During discharge, Li+ moves from anode to cathode

35. Voltage is 3.6/3.7V
Three times that of NiCad or NiMH
Much higher

36. The anode of lithium cell is made of reactive lithium metal.
Q.23
The anode of lithium cell is made of reactive lithium metal.
If the lithium anode is exposed to moisture and air, vigorous reactions will occur.
Thus, lithium ion cell is safer to use.

37. Continuous supply of fuel, H2

38. Continuous supply of oxygen
No need for recharging

39. Other fuels such as hydrocarbon, alcohol, or glucose are possible

40. Ni(s) and NiO(s) are catalysts for the half-cell reactions

41. Fuel Cells At anode: H2(g) + 2OH–(aq)  2H2O(l) + 2e– At cathode:
O2(g) + 2H2O(l) + 4e–  4OH–(aq)
Overall reaction:
2H2(g) + O2(g)  2H2O(l)

42. Q.24
Maximum energy that can be used to do useful work
= (2)(96485)(1.22) = 235 kJ mol1

43. Q.25
To increase the mobility of OH/K+ to balance the extra charges built up in half-cells.
[OH(aq)]  quickly at anode
[OH(aq)]  quickly at cathode
2. To increase the solubility of KOH

44. Q.26
Anode : -
CH4 + 2H2O CO2 + 8H+ + 8e
Cathode : -
2O2 + 8H+ + 8e H2O
Overall reaction : -
CH4 + 2O2  CO2 + 2H2O

45. Supplementary notes from HKDSE Chemistry

46. What is a fuel cell?

47. Fuel cell It is a primary cell.
It converts the chemical energy of a continuous supply of reactants (a fuel and an oxidant) into electrical energy.
The products are removed continuously.

48. How a fuel cell works
porous Ni electrodes
anode (−)
cathode (+)

49. How a fuel cell works Fuel : H2
hot KOH electrolyte ( 200°C) H2
hydrogen
Fuel : H2 e− e−
steam (exhaust)
Anode : H2(g) + 2OH(aq)  2H2O(g) + 2e

50. How a fuel cell works Fuel : H2 Oxidant : O2
hydrogen
Fuel : H2
Oxidant : O2 e− e−
steam (exhaust) H2 O2
oxygen
hydrogen
Cathode : O2(g) + 2H2O(g) + 4e  4OH(aq)

51. Functions of nickel electrodes:
act as electrical conductors that connect the fuel cell to the external circuit
act as a catalyst for the reactions

52. The reactions involved are:
At anode
H2(g) + 2OH–(aq)  2H2O(l) + 2e–
At cathode
O2(g) + 2H2O(l) + 4e–  4OH–(aq)
Overall reaction
2H2(g) + O2(g) ⇌ 2H2O(l)

53. Overall reaction
2H2(g) + O2(g) ⇌ 2H2O(l)
+ electrical energy
Direct reaction : -
Heat energy, light energy and sound energy (pop sound) will be released.
Other possible fuels include : ethanol, methanol, glucose solution…
But the cells have to be redesigned.

54. Applications of fuel cells
For remote locations, such as spacecraft, remote weather stations…
Continuous supply of fuel  No need to be replaced frequently
Fuel cells are used in space shuttle to provide electricity for routine operation.

55. high efficiency
e.g. hydrogen-oxygen fuel cells : 70%
much higher than internal combustion engines ( 20%) in motor cars.
Non-polluting
The only waste product of hydrogen-oxygen fuel cells is water. No greenhouse gases like CO2 or acidic gases like SO2 and NOx are emitted.
In fact, water vapor is a greenhouse gas due to its high specific heat capacity.

56. Fuel cells can also be used in electrical and hybrid vehicles.
A fuel cell car developed by DaimlerChrysler in Germany.

57. An MP3 player runs on methanol fuel cell in which methanol is used as fuel.
Fuel cells can be used in portable electronic products.

58. A portable fuel cell charger for mobile phones.
Fuel cells can be used in portable electronic products.

59. Application Features of fuel cells Examples
Power source
for remote
locations
high efficiency
high reliability
non-polluting
able to work continuously
spacecraft
remote weather stations
large parks
rural locations
The features of fuel cells and their applications.

60. Application Features of fuel cells Examples
Backup power
source
high reliability
non-polluting
able to work continuously
hospitals
hotels
office buildings
Transportation
quiet
high efficiency(70%)
electric vehicles
boats
But expensive
The features of fuel cells and their applications.

61. Application Features of fuel cells Examples
Portable
electronic
products
high efficiency
non-polluting
lightweight
can be refilled conveniently
notebook computers
mobile phones
MP3 players
handheld breathalyzers
The features of fuel cells and their applications.
Class practice 32.4

62. Class practice 32.4
The fuel cells used to power mobile phones and notebook computers are not hydrogen-oxygen fuel cells. Instead, they are called ‘direct methanol fuel cells (DMFC)’. The DMFC uses replaceable methanol cartridges for refilling. The fuel, methanol, is a liquid and can be fed directly in the cell for power generation.

63. At anode: CH3OH + H2O  6H+ + CO2 + 6e
Methanol and water react at the anode, producing H+. Positive ions (H+) are transported across the proton exchange membrane to the cathode where they react with oxygen to produce water. The products of the overall reaction are carbon dioxide and water.
Write the equations for the reactions at the anode and the cathode respectively. -2 +4
At anode: CH3OH + H2O  6H+ + CO2
+ 6e -2
At cathode: O2 + 4H  2H2O
+ 4e

64. (b) State one advantage of using methanol over hydrogen as fuel in the fuel cell.
Methanol is a liquid which is easier to handle than gaseous hydrogen during refilling. Or
Methanol poses a lower risk of explosion than hydrogen. (Any ONE)

65. (c) What are the potential dangers associated with using methanol fuel cells?
Methanol is flammable, if carelessly handled, it may catch fire.
Furthermore, methanol is a colourless liquid like water, yet it is highly poisonous.
If it is not stored or labelled properly, there is a danger of accidental poisoning.

66. Different types of fuel cells and their applications
The hydrogen-oxygen fuel cells discussed in Ch.32 is a type of Alkaline Fuel Cells (AFC).
The table below summarizes the main features of some fuel cells.

67. The summary of the main feature of some fuel cells.
Fuel cell type
Common electrolyte
Operating
temperature
System output
Electrical
efficiency
Applications
Proton ExchangeMembrane(PEMFC)
Solid organic polymer called poly-perfluoro-sulphonic acid
50–100°C
< 1 kW–250 kW
53–58% (transportation) 25–35% (stationary)
• Backup power
• Portable power
• Small distributed generation
• Transportation
Alkaline (AFC)
Aqueous solution of potassium hydroxide soaked in a matrix
below 80°C
10 kW–100 kW
60%
• Military applications
• Space projects
The summary of the main feature of some fuel cells.

68. The summary of the main feature of some fuel cells.
Fuel cell type
Common electrolyte
Operating
temperature
System output
Electrical
efficiency
Applications
Phosphoric Acid (PAFC)
Liquid phosphoricacid soaked in a matrix
150–200°C
50 kW– 1 MW (250 kW module typical)
> 40%
• Distributed generation
Molten Carbonate (MCFC)
Molten lithium, sodium, and / or potassium carbonates, soaked in a matrix
600–700°C
< 1 kW–
1 MW (250 kW module typical)
45–47%
• Electric utility
• Large distributed generation
The summary of the main feature of some fuel cells.

69. The summary of the main feature of some fuel cells.
Fuel cell type
Common electrolyte
Operating
temperature
System output
Electrical
efficiency
Applications
Solid Oxide (SOFC)
Solid zirconium oxide to which a small amount of yttrium(III) oxide is added
650–1000°C
< 1 kW–3 MW
35–43%
• Auxiliary power
• Electric utility
• Large distributed generation
The summary of the main feature of some fuel cells.

70. All of these fuel cells need fairly pure hydrogen gas as fuel.
A reformer is usually used in these fuel cells to generate hydrogen gas from liquid fuel like petrol except MCFC and SOFC. (refer to Example 34.1(c))
Class practice 34.1

71. Article reading
Read the article below and answer the questions that follow.
Microbial fuel cells−a greener and more
efficient source of electricity for tomorrow
Bacteria are very small (size ~1μm) organisms which can convert a huge variety of organic compounds into carbon dioxide, water and energy. The micro-organisms use the produced energy to grow and to maintain their metabolism.

72. However, by using a microbial fuel cell (MFC), we can
collect a part of this microbial energy in the form of electricity.
An MFC consists of an anode, a cathode, a proton or cation exchange membrane and an electrical circuit.

73. The general layout of an MFC.
anode
cathode
wastewater
glucose
H2O
(bacteria) H+ O2
CO2 e− H+
membrane
The general layout of an MFC.

74. The bacteria live in the anode compartment and convert a substrate such as glucose and wastewater into carbon dioxide, hydrogen ions and electrons. The electrons then flow through an electrical circuit to the cathode. The potential difference (Volt) between the anode and the cathode, together with the flow of electrons (Ampere) result in the generation of electrical power (Watt). The hydrogen ions flow through the proton or cation exchange membrane to the cathode. At the cathode, an electron acceptor is chemically reduced. Ideally, oxygen is reduced to water.

75. Microbial fuel cells have a number of potential uses
Microbial fuel cells have a number of potential uses. The first and most obvious is collecting the electricity produced for a power source. Virtually any organic material could be used to ‘feed’ the fuel cell. MFCs could be installed in wastewater treatment plants. MFCs are a very clean and efficient method of energy production.

76. Questions
Are microbial fuel cells (MFC) really fuel cells? Why?
Yes.
This is because a fuel (organic material) and an oxidant (oxygen) are used in MFC to generate electricity.

77. Questions
2. Why are microbial fuel cells (MFC) considered a greener source of energy?
Microbial fuel cells use wastewater as the source of fuel and produce CO2 and water which are harmless.
3. Suggest TWO substances that can be used as the ‘fuel’ for microbial fuel cells (MFC).
Glucose and wastewater

78. Overall reaction : 6(1) + (2)
Write balanced ionic equations for the reactions that occur at the cathode and the anode.
Cathode
O2(g) + H+(aq)  H2O(l) (1) 4
+ 4e 2
Anode
C6H12O6(aq)  CO2(g) H+(aq) (2)
+ 6H2O (l) 6 24
+ 24e
Overall reaction : 6(1) + (2)
C6H12O6(aq) + 6O2(g)  6CO2(g) + 6H2O(l)

79. Distributed generation
Class practice 34.1
One possible use of fuel cells with great potential of becoming more and more common is as ‘combined heat and power systems’ (CHP). A CHP is a small power station used to generate both electric power and heat energy for use in a block of flats, or in a factory.
Give three reasons to support the argument that ‘a fuel cell CHP is better than a diesel generator for use as a CHP.’
Distributed generation

81. 1. (1) A diesel generator has a lower efficiency than a fuel cell system. In other words, a diesel generator consumes more fuel to produce the same quantity of heat and electricity as compared to a fuel cell.
(2) A diesel generator causes pollution to the environment, producing smoke, bad smell, and a lot of NOx and SO2. A fuel cell system is clean and the exhaust is non-polluting, so it is more suitable for on-site energy production for a block of flats.

82. (3) A diesel generator is very noisy while a fuel cell operates quietly. This again is better for on-site power production.
(4) Renewable fuels such as glucose can be used in CHP while diesel used in diesel generator is non-renewable.

83. Phosphoric acid fuel cells (PAFC) are a suitable choice to be used in CHP. In this type of cells, the electrolyte used is liquid phosphoric acid soaked in a matrix.
(a) Write the ionic half equations at the cathode and the anode of PAFC respectively
At cathode: O2(g) + 4H+(aq) + 4e  2H2O(l)
At anode: 2H2(g)  4H+(aq) + 4e
(b) Write the overall equation for the cell reaction.
2H2(g) + O2(g)  2H2O(l)

84. Lithium-ion rechargeable batteries Lithium-ion polymer rechargeable
There are two types of rechargeable lithium cells:
Lithium-ion
rechargeable
batteries
Lithium-ion polymer
rechargeable
batteries

85. Lithium-ion rechargeable batteries
Lithium-ion rechargeable batteries are commonly used in portable electronic devices.

86. In a lithium-ion rechargeable battery, both the positive electrode and negative electrode contain lithium compounds.

87. Discharging Load electrons current negative positive electrode
separator
electrolyte

88. Charging Charger electrons current negative positive electrode
separator
Charger
electrolyte
current
electrons

89. a metal oxide fitted with Li+ ion
Positive electrode
e.g. Li1-xCoO2
a metal oxide fitted with Li+ ion
e.g. cobalt dioxide CoO2, manganese dioxide MnO2 or nickel dioxide NiO2
Negative electrode
graphite
lithium-carbon compound LixC6
Prevents electrolysis of water to give H2
Electrolyte
a lithium salt in an organic solvent

90. The chemical equations for the reactions are:
Positive electrode
discharging +4 +3
Li1-xCoO2 + xLi+ + xe− LiCoO2
charging
Negative electrode
discharging
charging
LixC C + xLi+ + xe−
It is the graphite in the lithium compound that loses electrons

91. Note that lithium ions themselves are neither oxidized nor reduced.
Overall reaction
Li1-xCoO2 + LixC LiCoO2 + 6C
discharging
charging
(+)
()
Note that lithium ions themselves are neither oxidized nor reduced.
The voltage of a lithium-ion rechargeable battery is 3.7 V.

92. Comparison with other cells
Feature
Comparison with other cells
High charge density
A lithium-ion rechargeable battery weighs about half that of a NiCd or NiMH cell of the same charge capacity.
High voltage (3.6–3.7 V)
A voltage range more suitable for many portable electronic devices like mobile phones, MP3 players, digital cameras, etc.
A summary of the comparison of lithium-ion rechargeable batteries with other cell types.

93. Comparison with other cells
Feature
Comparison with other cells
High drain capacity
Lithium-ion rechargeable batteries can discharge at much larger currents than NiCd or NiMH cells continuously for a longer period of time. This is very important for some applications such as the steady conversation over the mobile phones.
Environmentally preferred
Lithium-ion rechargeable batteries do not contain mercury, lead or cadmium, as the zinc-carbon cells, lead-acid accumulators or nickel-cadmium cells do.
A summary of the comparison of lithium-ion rechargeable batteries with other cell types.

94. Comparison with other cells
Feature
Comparison with other cells
No lithium metal
Lithium-ion rechargeable batteries contain lithium compounds instead of the reactive lithium metal. This makes lithium-ion rechargeable batteries safer for use and for transportation.
Long cycle life
Lithium-ion rechargeable batteries can be recharged and discharged for 1200 cycles within 3 years.
Low self-discharge rate
Lithium-ion rechargeable batteries only lose about 5% of the charge per month. NiCd and NiMH cells lose about 1% of the charge per day.

95. Comparison with other cells
Feature
Comparison with other cells
Fast charge possible
Lithium-ion rechargeable batteries can be fast charged to 70–80% of full capacity in one hour.
Wide range of operating temperatures
Lithium-ion rechargeable batteries can be discharged between the temperature range from –20°C to 60°C, and can be recharged between 0°C to 45°C.
A summary of the comparison of lithium-ion rechargeable batteries with other cell types.

96. Lithium-ion polymer rechargeable batteries (Li-poly / LiPo)
rechargeable batteries are now commonly used in mobile phones.

97. Advantages of Li-poly/ LiPo
The lithium-salt electrolyte is not held in an organic solvent as in the lithium-ion design, but in a solid polymer composite such as polyethene oxide or polyacrylonitrile.
Advantages of Li-poly/ LiPo
The battery can be made to any shape.
The rate of self-discharge is much lower compared with that of nickel-cadmium and nickel-metal hydride rechargeable batteries.
Class practice 34.2

98. Class practice 34.2
Lithium-ion rechargeable batteries use lithium compound instead of lithium metal as the anode. Explain why lithium metal should not be used in batteries.

99. Lithium metal, like other alkali metals (sodium, potassium, etc
Lithium metal, like other alkali metals (sodium, potassium, etc.) reacts vigorously with water to produce hydrogen and a corrosive, strongly alkaline solution LiOH.
2Li(s) + 2H2O(l)  2LiOH(aq) + H2(g)
If the seal of a cell with a lithium metal anode is broken, water or even moisture in the air may react with lithium, causing hydrogen and alkaline solution to leak out.
Hydrogen may cause explosion and the alkaline solution can cause severe skin burns.