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Primary Cells. A primary cell is a cell that can be used once only and cannot be recharged. The reactants cannot be regenerated. Primary cells. non-rechargeable. These cells are not rechargeable. ․ Zinc-carbon cells.

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slide1
Primary Cells

A primary cell is a cell that can be used once only and cannot be recharged.

The reactants cannot be regenerated

slide2
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.

slide3
Primary cells
  • These cells are not rechargeable.

․Alkaline manganese cells

slide4
Primary cells
  • These cells are not rechargeable.

․Silver oxide cells (button cells)

slide5
Primary cells
  • These cells are not rechargeable.

․Lithium primary cells (button cells)

slide7
At anode:

At cathode:

2MnO2(s) + 2NH4+(aq) + 2e– Mn2O3(s) + 2NH3(aq) + H2O(l)

Zn(s) Zn2+(aq) + 2e–

Zinc-carbon Cells

Overall reaction:

Zn(s) + 2MnO2(s) + 2NH4+(aq) Zn2+(aq) + Mn2O3(s) + 2NH3(aq) + H2O(l)

Ecell = +1.50 V

slide8
Zn(s) | Zn2+(aq) [2MnO2(s) + 2NH4+(aq)], [Mn2O3(s) + 2NH3(aq) + H2O(l)] | C(graphite)

Zinc-carbon Cells

The cell diagram for the zinc-carbon cell is:

Overall reaction:

Zn(s) + 2MnO2(s) + 2NH4+(aq) Zn2+(aq) + Mn2O3(s) + 2NH3(aq) + H2O(l)

slide9
At anode,

Zn(s) + 2OH(aq) ZnO(s) + H2O(l) + 2e

At cathode,

Ag2O(s) + H2O(l) + 2e 2Ag(s) + 2OH(aq)

Overall reaction :Zn(s) + Ag2O(s)  ZnO(s) + 2Ag(s)

slide10
At anode,

Zn(s) + 2OH(aq) ZnO(s) + H2O(l) + 2e

At cathode,

HgO(s) + H2O(l) + 2e Hg(l) + 2OH(aq)

Q.20(a)

HgO(s)

Overall reaction : Zn(s) + HgO(s)  ZnO(s) + Hg(l)

slide11
Q.20(b)

HgO(s)

= +0.098V – (1.216V) = 1.314V

slide12
Secondary Cells

Electrochemical cells that can be recharged.

Examples : -

Lead-acid accumulators

Nickel-cadmium cells (NiCad)

Nickel-Metal hydride(NiMH) cells

Lithium-ion cells

slide13
Lead grids coated with PbSO4(s)

Pb(s) + H2SO4(aq)  PbSO4(s) + H2(g)

slide14
Negative electrode :

PbSO4(s) + 2e Pb(s) + SO42(aq)

spongy lead

Lead grids coated with PbSO4(s)

During charging

slide15
Positive electrode :

PbSO4(s) + 2H2O(l)

PbO2(s) + 4H+(aq) + SO42(aq) + 2e

spongy PbO2

Lead grids coated with PbSO4(s)

During charging

slide16
Anode :

Pb(s) + SO42(aq) PbSO4(s)+ 2e

Lead grids coated with PbSO4(s)

During discharging

slide17
Cathode :

PbO2(s) + 4H+(aq) + SO42(aq) + 2e PbSO4(s) + 2H2O(l)

Lead grids coated with PbSO4(s)

During discharging

slide18
discharge

Pb(s) + PbO2(s) + 2H2SO4(aq) 2PbSO4(s) + 2H2O(l)

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)

Overall reaction : -

slide19
discharge

Pb(s) + PbO2(s) + 2H2SO4(aq) 2PbSO4(s) + 2H2O(l)

charge

Overall reaction : -

PbSO4 is coated on the electrodes,

The reversed processes are made possible.

slide20
discharge

Pb(s) + PbO2(s) + 2H2SO4(aq) 2PbSO4(s) + 2H2O(l)

charge

Overall reaction : -

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.

slide21
discharge

Pb(s) + PbO2(s) + 2H2SO4(aq) 2PbSO4(s) + 2H2O(l)

charge

Overall reaction : -

charge

Pb(s) and PbO(s) are on different electrodes

Direct reaction is not possible

Porous partition is not needed

slide22
discharge

Pb(s) + PbO2(s) + 2H2SO4(aq) 2PbSO4(s) + 2H2O(l)

charge

Overall reaction : -

During discharging, H2SO4 is being used up

The density of electrolyte solution 

The charging/discharging status can be monitored by a hydrometer.

slide23
Q.21

Ecell = Eocathode – Eoanode

= (1.69V) – (0.35V)=2.04V

slide24
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)

Nickel-cadmium cells – Nicad cells

Q.22(a)

slide25
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)

slide26
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.

slide27
On discharging,

Anode : -

MH(s) + OH(aq) M(s) + H2O(l) + e

+1

0

At cathode,

NiO(OH)(s) + H2O(l) + e

Ni(OH)2(s) + OH(aq)

+3

+2

Nickel metal hydride cell (NiMH)

slide28
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.

slide35
Voltage is 3.6/3.7V

Three times that of NiCad or NiMH

Much higher

slide36
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.

slide38
Continuous supply of oxygen

No need for recharging

slide41
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)
slide42
Q.24

Maximum energy that can be used to do useful work

= (2)(96485)(1.22) = 235 kJ mol1

slide43
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
slide44
Anode : -

CH4 + 2H2O CO2 + 8H+ + 8e

Cathode : -

2O2 + 8H+ + 8e 4H2O

Q.26

Overall reaction : -

CH4 + 2O2 CO2 + 2H2O

slide47
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.
slide48
anode (−)

cathode (+)

porous Ni electrodes

How a fuel cell works

slide49
hot KOH electrolyte ( 200°C)

Fuel : H2

steam (exhaust)

H2

hydrogen

How a fuel cell works

e−

e−

Anode : H2(g) + 2OH(aq)  2H2O(g) + 2e

slide50
Fuel : H2

steam (exhaust)

H2

O2

hydrogen

oxygen

How a fuel cell works

Oxidant : O2

e−

e−

H2

hydrogen

Cathode : O2(g) + 2H2O(g) + 4e  4OH(aq)

slide51
Functions of nickel electrodes:
  • act as electrical conductors that connect the fuel cell to the external circuit
  • act as a catalystfor the reactions
slide52
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)

slide53
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.

slide54
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.

slide55
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.

slide56
Fuel cells can also be used in electrical and hybrid vehicles.

A fuel cell car developed by DaimlerChrysler in Germany.

slide57
methanol
  • An MP3 player runs on methanol fuel cell in which methanol is used as fuel.

Fuel cells can be used in portable electronic products.

slide58
fuel cell charger
  • A portable fuel cell charger for mobile phones.

Fuel cells can be used in portable electronic products.

slide60
But expensive

The features of fuel cells and their applications.

slide61
Class practice 32.4

The features of fuel cells and their applications.

slide62
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.

slide63
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

0

-2

At cathode: O2 + 4H+2H2O

+ 4e

slide64
(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)

slide65
(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.

slide66
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.
slide70
Class practice 34.1
  • 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))
slide71
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.

slide72
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.

slide73
anode

cathode

wastewater

glucose

H2O

(bacteria)

H+

O2

CO2 e− H+

membrane

The general layout of an MFC.

slide74
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.
slide75
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.
slide76
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.

slide77
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

slide78
Write balanced ionic equations for the reactions that occur at the cathode and the anode.

Cathode

O2(g) + H+(aq)  H2O(l) (1)

4

2

+ 4e

Anode

+ 24e

C6H12O6(aq)  CO2(g) + H+(aq) (2)

+ 6H2O (l)

6

24

Overall reaction : 6(1) + (2)

C6H12O6(aq) + 6O2(g)  6CO2(g) + 6H2O(l)

slide79
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

slide81
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.

slide82
(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.

slide83
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)

slide84
There are two types of rechargeable lithium cells:

Lithium-ion

rechargeable

batteries

Lithium-ion polymer

rechargeable

batteries

slide85
Lithium-ion rechargeable batteries

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

slide86
In a lithium-ion rechargeable battery, both the positive electrode and negative electrode contain lithium compounds.
slide87
Load

electrons

current

negative

electrode

positive

electrode

separator

electrolyte

Discharging

slide88
Charger

electrons

current

positive

electrode

negative

electrode

separator

electrolyte

Charging

slide89
graphite

Prevents electrolysis of water to give H2

e.g. Li1-xCoO2

Positive electrode

a metal oxide fitted with Li+ ion

e.g. cobalt dioxide CoO2, manganese dioxide MnO2 or nickel dioxide NiO2

Negative electrode

lithium-carbon compound LixC6

Electrolyte

a lithium salt in an organic solvent

slide90
Li1-xCoO2 + xLi+ + xe− LiCoO2

discharging

LixC6 6C + xLi+ + xe−

charging

It is the graphite in the lithium compound that loses electrons

  • The chemical equations for the reactions are:

Positive electrode

discharging

+4

+3

charging

Negative electrode

slide91
discharging

Li1-xCoO2 + LixC6 LiCoO2 + 6C

charging

(+)

()

Overall reaction

  • Note that lithium ions themselves are neither oxidized nor reduced.
  • The voltage of a lithium-ion rechargeable battery is 3.7 V.
slide96
Lithium-ion polymer rechargeable batteries (Li-poly / LiPo)

Lithium-ion polymer

rechargeable batteries are now commonly used in mobile phones.

slide97
Class practice 34.2
  • 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.
slide98
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.

slide99
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.

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