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 Recharging is dangerous as it produces H2 and heat which results in an explosion.

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

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

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

Zinc-carbon Cells

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 = +1.50 V

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)

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)

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)

Q.20(b) HgO(s) = +0.098V – (1.216V) = 1.314V

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

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

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

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

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

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

Overall reaction : - Pb(s) + PbO2(s) + 2H2SO4(aq) 2PbSO4(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)

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

Overall reaction : - Pb(s) + PbO2(s) + 2H2SO4(aq) 2PbSO4(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.

Overall reaction : - Pb(s) and PbO(s) are on different electrodes Pb(s) + PbO2(s) + 2H2SO4(aq) 2PbSO4(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

Overall reaction : - During discharging, H2SO4 is being used up Pb(s) + PbO2(s) + 2H2SO4(aq) 2PbSO4(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.

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

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)

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)

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.

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

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.

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

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

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

During discharge, Li+ moves from anode to cathode

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

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.

Continuous supply of fuel, H2

Continuous supply of oxygen No need for recharging

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

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

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)

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

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

Q.26 Anode : - CH4 + 2H2O CO2 + 8H+ + 8e Cathode : - 2O2 + 8H+ + 8e 4H2O Overall reaction : - CH4 + 2O2  CO2 + 2H2O

Supplementary notes from HKDSE Chemistry

What is a fuel cell?

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.

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

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

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)

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

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)

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.

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.

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.

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

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

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

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.

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.

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

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.

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

(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)

(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.

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.

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.

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.

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.

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

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.

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.

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

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.

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.

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.

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

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)

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

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.

(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.

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)

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

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

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

Discharging Load electrons current negative positive electrode separator electrolyte

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

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

The chemical equations for the reactions are: Positive electrode discharging +4 +3 Li1-xCoO2 + xLi+ + xe− LiCoO2 charging Negative electrode discharging charging LixC6 6C + xLi+ + xe− It is the graphite in the lithium compound that loses electrons

Note that lithium ions themselves are neither oxidized nor reduced. Overall reaction Li1-xCoO2 + LixC6 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.

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.

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.

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.

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.

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

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

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.

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.