Form 5 Chemistry Notes Chapter 3 Oxidation and Reduction

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FORM 5 CHEMISTRY
CHAPTER 3 OXIDATION AND REDUCTION
3.1 Redox Reactions
Oxidation and Reduction Reactions
 Oxidation can be defined as
- Acceptance (Gain) of oxygen
- Donation (Loss) of hydrogen
- Loss of electrons
- Increase in oxidation number of element
 Reduction can be defined as
- Loss of oxygen
- Gain of hydrogen
- Gain of electrons
- Decrease in oxidation number of element
Oxidation in Terms of Gain of Oxygen
 OXIDATION: Chemical reaction in which oxygen is added to a substance.
 If substance gains oxygen during reaction, it is said to be oxidised.
 When calcium burns in oxygen, calcium is oxidised because it gains oxygen in this reaction.
2Ca + O2  2CaO
Oxidation in Terms of Loss of Hydrogen
 OXIDATION: Loss of hydrogen from a substance.
 If substance loses hydrogen during a reaction, it is said to be oxidised.
 When hydrogen sulphide gas is mixed with chlorine gas at room temperature, yellow precipitate
of sulphur is formed and hydrogen chloride gas is released. Hydrogen sulphide loses oxygen and
is oxidised to sulphur.
H2S + Cl2  2HCl + S
Reduction in Terms of Loss of Oxygen
 REDUCTION: Loss of oxygen from a substance.
 If a substance loses oxygen during reaction, it is said to be reduced.
 When a mixture of zinc powder and copper(II) oxide is heated, copper(II) oxide lost its oxygen. It
is said to be reduced to metallic copper.
Zn + CuO  ZnO + Cu
Reduction in Terms of Gain of Hydrogen
 REDUCTION: Addition of hydrogen to a substance.
 If a substance gains hydrogen during reaction, it is said to be reduced.
 When a mixture of hydrogen and chlorine is exposed to sunlight, vigorous reaction occurs and
white fumes of hydrogen chloride are produced. In this reaction, chlorine has gained hydrogen,
chlorine has been reduced.
H2 + Cl2  2HCl

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Redox Reactions
 Oxidation and reduction always take place together.
 REDOX REACTION: Reaction in which both oxidation and reduction take place simultaneously.
 In redox reaction, when one substance in reaction is oxidised, the other substance is reduced
 When steam is passed over heated magnesium, magnesium oxide and hydrogen are produced.
Magnesium is oxidised whereas water is reduced.
Mg + H2O  MgO + H2
 Respiration is a redox process. Food is oxidised and oxygen molecules accept electrons and are
reduced to water.
C6H12O6 + 6O2  6CO2 + 6H2O
 In contrast, photosynthesis is also a redox reaction. Electrons are removed from water molecules
and are used to reduce carbon dioxide to sugar.
6CO2 + 6H2O  C6H12O6 + 6O2
Oxidising and Reducing Agents
 OXIDISING AGENT: Substance that brings about oxidation in another substance and is itself
reduced
 REDUCING AGENT: Substance that bring about reduction in another substance and is itself
oxidised.
Oxidising agents Reducing agents
 Chlorine and bromine
 Acidified potassium manganate (VII)
 Acidified potassium dichromate (VI)
 Concentrated nitric acid
 Metals such as sodium, magnesium, zinc and
aluminium
 Sulphur dioxide gas and hydrogen sulphide gas
 Sodium sulphite and sodium thiosulphate
 Potassium iodide
Reaction Oxidising agent Reducing agent
Reaction between copper(II) oxide and
carbon:
2CuO + C  2Cu + CO2
Copper(II) oxide
 Copper(II) oxidises
carbon to carbon
dioxide
 Reduced to copper
Carbon
 Reduces copper(II)
oxide to copper.
 Oxidised to carbon
dioxide

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Oxidation Number
 OXIDATION NUMBER / OXIDATION STATE: Arbitrary charge assigned to element according to
a set of rules.
 Ionic compounds
- Ionic compound can contain monoatomic ions (E.g.: Na+, Cl-) or polyatomic ions (E.g.: NH4
+ or
SO4
2-)
- For monoatomic ion, oxidation number is charge on ions.
- MgO is ionic compound In MgO, Mg exists as Mg2+ and oxygen exists as O2-. Mg is said to
have oxidation number of +2 and oxygen has oxidation number of –2.
 Covalent compounds
- CO2 is covalent compound. However, when determining oxidation number of C and O, the
molecule need to be considered exists as ions.
- Each O atom is considered an oxide (O2-) and carries charge of –4.
- Each carbon ion carries charge of +4, so that CO2 exists as neutral molecule.
Rules for Assigning Oxidation Number
 Atom or molecule in free state has oxidation number of 0.
Element Formula Oxidation number
Hydrogen H2 0
Chlorine Cl2 0
Sulphur S 0
 For monoatomic ions, oxidation number equals to charge in ion.
Simple ion Formula of ion Oxidation number
Hydrogen ion H+ +1
Magnesium ion Mg2+ +2
Nitride ion N3- –3
 Sum of oxidation states of all atoms present in formula of compound is 0.
- CaCO3 [Ca: +2; C: +4, O: –2; (+2) + (+4) + 3(–2) = 0]
 For polyatomic ion, sum of oxidation numbers of all atoms equals the charge on ion.
- SO4
2- [S: +6; O: –2; (+6) + 4(–2) = –2]
 Oxidation number of fluorine (–1) remains unchanged in all compounds.
- F2O (Oxidation number of O is +2)
- BrF3 (Oxidation number of Br is +3)
 Chlorine, bromine and iodine usually have oxidation number of –1 except when combined with a
more electronegative element.
- HClO (H: +1; C: +1; O: –1)
- KIO3 (K: +1; I: +5; O: –2)

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 Oxidation number of hydrogen in all its compounds is +1 except in metal hydrides (–1).
- NaH (Na: –1; H: +1)
- CaH2 (Ca: +2; H: –1)
 Oxidation number of oxygen in all its compounds is –2 except in fluorine compound (stated above)
and peroxides.
- H2O2 (H: +1; O: –1)
 Metals usually have positive oxidation number. (Group 1: +1; Group 2: +2)
 Some metals show different oxidation in their compounds.
 Non-meals usually have negative oxidation numbers
Compound MnSO4 MnO2 K2MnO4 KMnO4
Oxidation number of
manganese
+2 +4 +6 +7
- Cl, Br and I can have positive or negative oxidation number depending on elements which
combine to them.
Chlorine compound HCl HClO HClO2 ClO2 HClO3 HClO4
Oxidation number of
chlorine
–1 +1 +3 +4 +5 +7
Nitrogen compound NH3 N2O NO NO2
- NO2 NO3
-
Oxidation number of
nitrogen
–3 +1 +2 +3 +4 +5
IUPAC Nomenclature of Inorganic Compounds
 IUPAC system is used to name inorganic compounds in order to avoid confusion that may arise due
to elements having different oxidation number.
 For example, there are two oxides of copper, Cu2O and CuO. Cu2O is brown powered whereas
CuO is black powder. Roman numerical figures (I) and (II) refer to oxidation numbers of copper in
compound.

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Oxidation Number and IUPAC Nomenclature
 For compound that contains metal with more than one oxidation number, Roman numerical figure is
stated in brackets after name of metal to show oxidation number of metal.
- For example Sn forms two types of chlorides, SnCl2 (ionic) named tin(II) chloride and SnCl4
(covalent) named tin(IV) chloride.
Formula of compound Oxidation number of metal IUPAC name
FeCl2 +2 Iron(II) chloride
Mn(NO3)2 +2 Manganese(II) nitrate
MnO2 +4 Manganese(IV) oxide
 Metallic elements in Groups 1, 2 and 3 of Periodic Table always have oxidation number +1, +2
and +3 respectively. Roman numerical figure is not used in naming compound if metal shows only
one oxidation state.
Formula of compound Oxidation number of metal IUPAC name
K2SO4 +1 Potassium sulphate
Mg(NO3)2 +2 Magnesium nitrate
AlCl3 +3 Aluminium chloride
 For negative ion that contains metal with more than one oxidation state, Roman number is stated
in brackets after name o metals and name of metal ends with -ate.
- Manganate(VII) refers to negative ion containing manganese metal with oxidation number +7
which is MnO4
-
- Chromate(VI) refers to negative ion containing chromium metal with oxidation number + 6
which is CO4
2-
- Hexacyanoferrate(III) refers to negative ion containing six cyano (CN-) groups and iron metal
with oxidation number +3 which is [Fe(CN)6]3+
Formula of compound Oxidation number of metal IUPAC name
K2MnO4 +6 Potassium manganate(VI)
K3Fe(CN)6 +2 Potassium hexacyanoferrate(II)
K4Fe(CN)6 +3
Potassium
hexacyanoferrate(III)

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 OXOANIONS: Anions that consist of an oxygen atom and another non-metallic atom.
- E.g.: NO3
-, SO4
2-
- For non-metal that shows more than one oxidation number in its oxoanion, Roman number
stated in brackets refers to oxidation number of non-metal.
Molecular formula of
compound
Oxidation number IUPAC name
Common name of
compound
Na2SO3 +4 Sodium sulphate(IV) Sodium sulphite
Na2SO4 +6 Sodium sulphate(VI) Sodium sulphate
NaNO2 +3 Sodium nitrate(III) Sodium nitrite
NaNO3 +5 Sodium nitrate(V) Sodium nitrate
HNO2 +3 Nitric(III) acid Nitrous acid
HNO3 +5 Nitric(V) acid Nitric acid
H2SO4 +6 Sulphuric(VI) acid Sulphuric acid
Oxidation and Reduction in Terms of Changes in Oxidation Numbers
 Most redox reactions occur without involving hydrogen or oxygen, they are discussed in terms of:
- Changes in oxidation numbers
- Transfer (Gain / Loss) of electrons
 OXIDATION: Process in which oxidation number of element is increased.
 REDUCTION: Process in which oxidation number of
element is increased.
 Iron metal oxidised to iron(III) chloride as its oxidation number increases from 0 to 3.
 Chlorine is reduced to chloride ion because its oxidation number decreases from 0 to –1.
 A reaction is not a redox reaction if the substances involved in reaction do not undergo an
changes in oxidation numbers:
- Reaction between NaOH and H2SO4 is neutralization reaction and not redox reaction.
Oxidation numbers of all elements (Na, O, H and S) are same before and after the reaction.

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Oxidation and Reduction in Terms of Electron Transfer
 In terms of electron transfer:
- OXIDATION: Loss of electrons from a substance.
- REDUCTION: Gain of electron by a substance.
 Oxidising agents are electron acceptors while reducing agents are electron donors.
 If a coil of Cu is placed in solution of AgNO3, Cu slowly dissolves and solution turns blue. Copper
coil becomes coated with layer of Ag metal.
- Overall equation: Cu + 2AgNO3  Cu(NO3)2 + 2Ag
- Ionic equation: Cu + 2Ag+  Cu2+ + 2Ag
- Each Ag+ ion accepts one electron to form Ag atom (reduction)
- Each Cu atom donates two electrons and are converted to Cu2+ ion in aqueous solution
(oxidation)
 Combustion of metals in chlorine
- When hot copper foil is placed in gas jar of chlorine, vigorous reaction occurs and green
precipitate of copper(II) chloride, CuCl2 is formed.
- Cu + Cl2  CuCl2
 Cu  Cu2+ + 2e-
 Loses electrons
 Undergoes oxidation
 Oxidised to copper(II) ion, Cu2+
 Acts as reducing agent
 Cl2 + 2e-  2Cl-
 Gains electrons
 Undergoes reduction
 Reduced to chloride ion, Cl-
 Acts as oxidising agent

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 Combustion of metals in oxygen
- When metal burn in oxygen
 Metals undergo oxidation by losing electrons to form metal ions
 Oxygen undergoes reduction by gaining electrons to form oxide ions (O2-)
- 2Mg + O2  2MgO
 Mg  Mg2+ + 2e-
 Loses electrons
 Undergoes oxidation
 Oxidised to magnesium ion, Mg2+
 Acts as reducing agent
 O2 + 4e-  2O2-
 Gains electrons
 Undergoes reduction
 Reduced to oxide ion, O2-
 Acts as oxidising agent

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Oxidation of Fe2+ to Fe3+
 Potassium manganate(VII) can oxidise Fe2+ to Fe3+.
- When acidified KMnO4 solution is added to solution of Fe2+ salt, decolourisation occurs. MnO4
ions are reduced to Mn2+ while Fe2+ ions are oxidised to Fe3+
- MnO4
- + 8H+ + 5Fe2+  Mn2+ + 4H2O + 5Fe3+
 Fe2+  Fe3+ + e-
 MnO4
- + 8H+ +5e-  Mn2+ + 4H2O
- Formation of Fe3+ ions can be confirmed by using sodium hydroxide solution. When sodium
hydroxide solution is added to reaction product, brown precipitate of Fe(OH)3 insoluble in
excess NaOH is obtained. (Fe3+ + 3NaOH  Fe(OH)3 + 3Na+)
 Other oxidising agents
- Chorine gas / Chlorine water
 Cl2 + 2Fe2+  2Fe3+ + 2Cl-
- Liquid bromine
 Br2+ 2Fe2+  2Fe3+ + 2Br-
- Acidified potassium dichromate(VI) solution
 Cr2O7
2- + 14H+ + 6Fe2+  6Fe3+ + 2Cr3+ + 7H2O
- Concentrated nitric acid
 HNO3 + 3Fe2+ + 3H+  3Fe3+ + 2H2O + NO
- Acidified hydrogen peroxide
 H2O2 + 2H+ + 2Fe2+ --. 2Fe3+ + 2H2O
Reduction of Fe3+ to Fe2+
 When Na2SO3 solution is added to FeCl3, and the mixture is acidified with dilute H2SO4, colour of
solution changes rom brown to light green.
- SO3
2- + H2O + 2Fe3+  2Fe2+ + H2SO4
- Na2SO3 acts as reducing agents and reduced Fe3+ ions to Fe2+ ions and is itself oxidised to
SO4
2-
 Fe3+ + e-  Fe2+
 SO3
2- + H2O  SO4
2- + 2H+ + 2e-
- Formation of Fe2+ can be confirmed by using NaOH solution. When NaOH solution is added to
reaction product, dirty green precipitate of Fe(OH)2, insoluble in excess NaOH is obtained.
(Fe2+ + 2NaOH  Fe(OH)2 + 2Na+)
 Other reducing agents
- Metals more electropositive than iron (E.g.: Zinc)
 Zn + Fe3+  2Fe2+ + Zn2+
- Sulphur dioxide
 SO2 + 2H2O + 2Fe3+  2Fe2+ + 2H+ + H2SO4
- Potassium iodide
 2KI + 2Fe3+  2Fe2++ 2K+ + I2
- Hydrogen sulphide
 H2S + 2Fe3+  2Fe2+ + 2H+ + S
- Tin(II) chloride solution
 Sn2+ + 2Fe3+  2Fe2+ + Sn4+

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Displacement of Metals from Their Salt Solutions
 The higher the position of metal in electrochemical series
- The more electropositive the metal
- The more readily the metal donates electrons to form positive ions
- The more easily the metal will undergo oxidation
 Electropositive metals are strong reducing agents. In contrast, metallic ions of electropositive
metals are weak oxidising agents.
- The strength of metal as reducing agent increases on going up electrochemical series
- The strength of metallic ion as an oxidising agent increases on going down the series.
 Consider the formation of Na+ from Na
- Na metal is placed at high position in electrochemical series. Na metal donates electrons very
easily.
- Conversely, Na+ have weak tendency to accept electrons. Since oxidising agents are electron
acceptors Na+ ions are weak oxidising agents.
 DISPLACEMENT REACTION: Reaction in which one element displaces another element from its
salt solution.
- The more electropositive metal will displace a less electropositive metal from salt solutions of
less electropositive metal.
 Transfer of electrons occurs during displacement reactions.
- The more electropositive metal donates electrons and acts as reducing agent. Metal
undergoes oxidation and is oxidised to its metal ions.
- Metal ion in aqueous solution acts as oxidising agent. Metal ions undergo reduction and are
reduced to its metal.
Displacement of Copper by Zinc from Copper(II) Sulphate Solution
 Reaction between copper(II) sulphate and zinc
- Zn + CuSO4 + ZnSO4 + Cu
- Zinc is more electropositive than copper, it displaces copper from its salt.
 Displacement reaction is a redox reaction
- Zn  Zn2+ + 2e-
- Cu2+ + 2e-  Cu
 When Cu2+ is displaced, concentration of Cu2+ ions in solution decreases. This causes blue colour to
fade.
Displacement of Silver by Copper from Silver Nitrate Solution
 Reaction between copper and silver nitrate
- Cu + 2AgNO3  Cu)NO3)2 + 2Ag
- Copper is more electropositive than silver. It displaces silver from its salt.
 Redox reaction
- Cu  Cu2++ + 2e-
- Ag+ + e-  Ag
 When copper dissolves in silver nitrate solution, formation of copper(II) ion causes solution to turn
blue. Intensity of blue colour increases as more copper is dissolved

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Displacement of Halogens from Halide Solutions
 Reactivity of halogens can be used to predict whether displacement reactions o halogens can
occur or not
 More reactive halogen will displace less reactive halogen from the solution of its halide ions.
Cl2 + 2KBr  2KCl + Br2
 Chlorine displaces bromine from aqueous solution of bromide ions. But bromine cannot displace
chlorine from aqueous solution of chloride ions.
Br2 + KCl  No reaction
 Colours of halogen
Halogen
Concentrated aqueous
solution
Dilute aqueous solution
Chlorine Greenish-yellow Colourless
Bromine Brown Yellow
Iodine Brown Yellow
 Halogens can be identified by adding 1, 1, 1-trichloroethane (CH3CCl3) to its aqueous solution.
Water and 1, 1, 1-trichloroethane are immiscible and two layers are formed.
- Upper layer is water and lower layer is 1, 1, 1-trichloroethane
- Colours of 1, 1, 1-trichloroethane
Halogen Colour
Chlorine Colourless
Bromine Brown
Iodine Purple

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Redox Reactions by the Transfer of Electrons at a Distance
 If solution containing oxidising agent is separated from solution containing reducing agent by
electrolyte, redox reaction can still occur by transfer of electrons at a distance.
 Electrons are transferred by connecting wire in external circuit, from reducing agent to oxidising
agent.
 Reducing agent acts as negative terminal (loss of electrons).
 Oxidising agent acts as positive terminal (gain of electrons).
 Sulphuric acid acts as salt bridge.
- Separate oxidising agent from reducing agent
- Complete electric circuit so that ions can move through it
 NaCl, KCl, NaNO3 and KNO3 can also be used as salt bridges.
Reaction between Potassium Iodide and Acidified Potassium Manganate(VII) by Transfer of Electrons
at a Distance
 Deflection of galvanometer needle shows that electrons flow in external circuit from carbon
electrode immersed in KI solution (negative electrode) to carbon electrode immersed in KMnO4
solution (positive electrode)
 Changes at negative electrode
- Colourless layer of KI slowly changes to yellow
- Oxidation reaction occurs at negative e electrode: 2I-  I2 + 2e-
- Electrons released during oxidation then flow through connecting wire (external circuit) to
positive electrode and are accepted by MnO4
- ions. Hence, MnO4
- ions acts as oxidising
agent.

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 Changes at positive electrode
- Purple layer of KMnO4 slowly becomes colourless
- Reduction reaction occurs at positive electrode: MnO4
- + 8H+ + 5e-  Mn2+ + 4H2O
 Overall reaction
- 2MnO4
- + 10I- + 16H+  2Mn2+ + 5I2 + 8H2O
3.2 Rusting as a Redox Reaction
Conditions for the Rusting of Iron
 RUSTING: Redox reaction between iron, oxygen and water to form brown substance called rust.
 Rust is hydrated iron(III) oxide, Fe2O3.xH2O. Composition of water is not constant.
 Conditions required for rusting:
- Presence of air
- Presence of water
Corrosion of Metals
 Corrosion of metals is redox reaction in which a metal is oxidised spontaneously at room
temperature with release of electrons to form metal ions.
 Most metals corrode readily in air. When corrosion occurs, metal surface loses its shine and
becomes dull.
- Metals react slowly in oxygen in air to form metal oxides on metal surface.
 Metals that are more reactive will corrode more readily.

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 The higher the position of metal in reactivity series, the easier it is for the metal to donate its
electrons and be corroded.
 Layer of Al2O3 can be made thicker by electrolysis (anodising).
- Anodising is used to protect aluminium rom rusting.
- Electrolytic process using aluminium as anode.
- During electrolysis, a layer of aluminium oxide is deposited on surface of aluminium.
Rusting in Terms of Oxidation and Reduction
 Rusting of iron is an electrochemical process that occurs spontaneously. When iron is in contact with
water, simple chemical cell is formed.
 Consider a drop of water on metal surface:
- Centre of drop of water
 Area where there is a lack of oxygen
 Act as negative terminal (anode)
- Side of drop of water
 Area where it is rich in oxygen
 Act as positive terminal (cathode)

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 Stages in rusting of iron
- In the centre of water droplet (anode)
 Iron rusts via oxidation process to form iron(II) ions
 Fe  Fe2+ + 2e-
 Electrons flow to the edge of water droplet through iron surface.
- At edge of water droplet
 Oxygen accepts electrons from oxidation of iron and is reduced to hydroxide ions.
 O2 + 2H2O + 4e-  4OH-
- Formation of Fe(OH)2
 Fe2+ + 2OH-  Fe(OH)2
- Formation of rust
 Iron(II) hydroxide produced is oxidised by oxygen to form iron(III) hydroxide, which then
decomposes to hydrated iron(III) oxide.
 4Fe(OH)2 + 2H2O + O2  4Fe(OH)3
 Fe(OH)3  Fe2O3.xH2O
 Equation for the redox reaction: 2Fe + O2 + 2H2O  2Fe(OH)2
 Overall equation for rusting of iron: 4Fe + 3O2 + 2xH2O  2Fe2O3.xH2O
 Rate of rusting of iron is increased if strong electrolyte is present. Rusting of iron occurs rapidly in
areas near sea or industrial areas.
- Sea air contains salts such as NaCl and MgCl.
- In industrial area, air is polluted by acidic gases such as SO2 and NO2.
- These substances increase he electrical conductivity of water, thus making water a better
electrolyte.
 Besides corrosion of iron and steel, corrosion of other metals also can occur. The main causes of
corrosion of metals are attack by chemicals such as acids, damp air or electrochemical corrosion.
Prevention of Rusting of Iron
 Rusting of iron can be prevented if iron is in contact with more electropositive metal.
 Conversely, rate of rusting of iron is increased if iron is in contact with less electropositive metal.
 When two metals are in contact, the greater the difference in electropositivity between these two
metals, the faster the more electropositive metal will rust.

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Methods Used for the Prevention of Rusting
 Using a protective layer
- Rusting of iron and steel can be prevented by keeping them away from air and water.
- A layer of paint, oil, grease or plastic coating protects iron surface from coming into contact
with air and water.
- Without presence of both air and water, rusting of iron cannot occur.
 Using less electropositive metals
- Plating iron with tin
 Tin is not an electropositive metal and is resistant to oxidation by water and air.
 Tin plating makes article shiny and more attractive in appearance
 Disadvantage: Tin is less electropositive than iron, if tin coating is broken, iron beneath will
rust even more rapidly.
- Plating iron with chromium
 Chromium is metal that is resistant to rusting. When chromium is exposed to water and air,
impermeable, non-brittle oxide layer is formed.
 Oxide layer acts as protective layer to prevent iron beneath it from coming into contact
with water and air in atmosphere.
 Using more (electropositive) reactive meals
- GALVANISING: Coating of iron or steel with zinc for protection from corrosion.
- Carried out by dipping iron object into molten zinc or by electroplating.
- Zinc-coated iron is known as galvanized iron.
 Even if layer of zinc is scratched, iron beneath it does not rust.
 Zinc is more electropositive than iron and will corrode first.
 This method is called cathodic protection. Metal zinc is known as sacrificial metal because
zinc is “sacrificed” in protection of iron from rusting.
 Rusting in ships is prevented by fixing zinc bars to part of the ship submerged in water
 Rusting in underground iron pipes is prevented by having blocks of magnesium attached to
iron pipes.
- Magnesium is more electropositive than iron

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 Using alloys
- Best known rust-resistant alloy of iron is stainless steel.
- Contains 10 – 20% nickel and 10 – 25% chromium. When exposed to air, hard layer of
chromium(III) oxide is formed on surface of iron and prevents iron from rusting.
- Stainless steel is used to make surgical instruments and kitchen wares such as knives, forks and
spoons.
3.3 The Reactivity Series of Metals and Its Applications
Reactivity of Metal with Oxygen
 Most metals form metal oxides when heated or burnt in air.
 Different metals have different reactivity with oxygen.
 Reactivity of metals with oxygen can be compared by observing the flame of glow produced
when metal is heated in oxygen.
 The more reactive the metal is with oxygen, the more brightly and rapidly the metal burns.
 Oxygen used for burning metals is supplied by heating KMnO4, KNO3 or mixture of KClO3 and
MnO2.

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Reactivity Series of Metals
 Reactivity series of metals that do not include hydrogen and carbon.
 Position of metal in reactivity series can also be determined by reaction between metal and oxide
of another metal.
 If metal X is more reactive than metal Y with oxygen, metal X will displace metal Y from its oxide
when a mixture of powdered X and oxide of metal Y is heated.
X + oxide of Y  oxide of X + Y
- Mg + CuO  MgO + Cu
- Cu + MgO  No reaction
 The displacement reaction can be considered in terms of oxidation and reduction.
- Zn + PbO  ZnO + Pb
- In this reaction between Zn and PbO, Zn acts as reducing agent and reduced PbO to lead.
- Conversely, PbO acts as oxidising agent and oxidises Zn to ZnO.
Heating Carbon with Metal Oxides
 Position of carbon in reactivity series can be determined by heating carbon with metal oxides.
 When mixture of carbon and oxide of metal X is heated strongly, reaction will occur if carbon is
more reactive than metal X.
 Carbon acts as reducing agent and oxide of metal X is reduced to metal X.
 Carbon + oxide of metal X  metal X + carbon dioxide
 Conversely, if carbon does not remove oxygen, means that carbon is less reactive than the metal
in oxide.

19. 여철우-화학|19
Heating Carbon Dioxide with Metals
 Ability of metal to remove oxygen from carbon dioxide can be used to determine the position of
carbon in reactivity series.
 Na, K, Ca, Mg and Al are more reactive than carbon. These metals will therefore react with
carbon dioxide and remove oxygen from carbon dioxide
 Metal + CO2  metal oxide + C
 When a piece of burning Mg ribbon is put into CO2 in as jar, Mg will continue to burn for a short
time. Black specks of carbon can be seen on sides of gas jar and Mg burns to form white powder.
(MgO)
- 2Mg + CO2  2MgO + C
 This reaction shows that magnesium
- is more reactive than carbon with oxygen
- acts as reducing agent
- reduces carbon dioxide to carbon
- is itself oxidised to MgO
 The higher the element is in the reactivity series, the stronger it acts as reducing agent in redox
reaction.
 Conversely, if metal does not remove oxygen from CO2, it implies that metal is less reactive than
carbon.
The Position of Hydrogen in Reactivity Series of Metals
 Position of hydrogen in reactivity series can be determined by passing dry hydrogen gas over hot
metal oxides.
 If hydrogen is more reactive than metal X, hydrogen will reduce the oxide of metal X to metal X.
Hydrogen + oxide of metal X  H2O + metal X
 If reaction between hydrogen gas and metal oxide occurs, a flame will spread throughout the
metal oxide and metal is produced.
 If hydrogen gas does not remove oxygen from metal oxide, hydrogen is less reactive with oxygen
than metal.
 Hydrogen used for reducing metal oxides to metals can be produced from reaction between
dilute H2SO4 or dilute HCl and Zn.
Zn + H2SO4  ZnSO4 + H2
Zn + 2HCl  ZnCl2 + H2
 H2 gas is dried by passing through drying agent such as concentrated H2SO4 or anhydrous CaCl2.

20. 20|여철우-화학
Position of Carbon and Hydrogen in Reactivity Series
 Reactivity series that includes other carbon and hydrogen.
 Reactions between metal oxides and carbon // metal oxides and hydrogen
Metal oxide Reaction with carbon Reaction with hydrogen
Potassium oxide (K2O)
Sodium oxide (Na2O)
Calcium oxide (CaO)
Magnesium oxide (MgO)
Aluminium oxide (Al2O3)
No reaction No reaction
Zinc oxide (ZnO) ZnO + C  2Zn + CO2 No reaction
Iron(III) oxide (Fe2O3)
Tin(IV) oxide (SnO2)
Lead(II) oxide (PbO)
Copper(II) oxide (CuO)
Silver oxide (Ag2O)
2Fe2O3 + 3C  4Fe + 3CO2
SnO2 + C  S + CO2
2PbO + C  2Pb + CO2
2CuO + C  2Cu + CO2
2Ag2O + C  4Ag+ CO2
Fe2O3 + 3H2  2Fe +
3H2O
SnO2 + 2H2  Sn + 2H2O
PbO + H2 --. Pb + H2O
CuO + H2  Cu + H2O
Ag2O + H2  2Ag + H2O
Extraction of Metals from Their Ores
 Most metals in metal ores exist in forms of oxides, carbonates and sulphides in Earth‟s crust.
 Extraction of metals involves reduction of metal ores to metals.
 Two main methods:
- Electrolysis of metal compounds in molten state (Metals higher than carbon in reactivity series)
- Reduction of metal oxides by carbon (Metals lower than carbon in reactivity series)
Metal Method of extraction
K, Na, Ca, Mg Electrolysis of metal chlorides in molten state
Al Electrolysis of Al2O3 in molten state
Zn, Fe, Sn, Pb Heating metal oxides with carbon
Cu, Hg Heating metal sulphides in air
Ag, Au Exists as free elements in Earth‟s crust

21. 여철우-화학|21
Extraction of Iron from Its Ore
 Important iron ores are haematite and magnetite.
- Haematite contains iron(III) oxide, Fe2O3
- Magnetite contains triiron tetroxide, Fe3O4
 Extraction of iron from haematite or magnetite is carried out in blast furnace by reduction using
carbon
- Raw materials required
 Mixture of iron ore, coke (carbon) and limestone (CaCO3) is put in blast furnace Hot air is
blown into furnace from bottom.
- Production of carbon dioxide
 In lower section of blast furnace, oxygen in hot air reacts with coke to from carbon
dioxide
C +CO2  CO2
 At high temperatures, limestone decomposes into quicklime (CaO) and carbon dioxide
CaCO3  CaO + CO2
- Production of carbon monoxide
 In upper section of blast furnace, carbon dioxide reacts with coke to produce carbon
monoxide
C + CO2  2CO

22. 22|여철우-화학
- Reduction of iron ore to iron
 In upper section of blast furnace, where temperature is about 400 – 800 °C, iron ore is
reduced by carbon monoxide to iron
Fe2O3 + 3CO  2Fe + 3CO2
Fe3O4 + 4 CO  3Fe + 4CO2
 In lower section of blast furnace, iron ore is reduced by coke to iron
Fe2O3 + 3C  2Fe + 3CO
Fe3O4 + 2C  3Fe + 2CO2
 In these reactions, carbon and carbon monoxide act as reducing agent.
 Molten iron produced flows to bottom of blast furnace and is collected. Molten iron is
poured into moulds and set aside to solidify.
 Removal of impurities
- In blast furnace, CaO is produced from decomposition of CaCO3. It then reacts with silica to
form slag.
CaO + SiO2  CaSiO3
- Molten slag floats on top of iron. Lag and iron are separated through tap at bottom of
furnace.
- CaO acts as basic oxide, whereas SiO2 acts as acidic oxide.
- CaSiO3 produced during extraction of iron is used mainly for road surfacing.
 Summary of extraction of iron
 Iron produced is not pure iron and contains about 5% carbon. This is called cast iron.

23. 여철우-화학|23
Thermite Process
 THERMITE PROCESS: Displacement reaction between aluminium and iron(III) oxide to produce
iron.
2Al + Fe2O3  Al2O3 + 2Fe
 Magnesium acts as fuse to ignite mixture.
 When mixture of magnesium powder and barium peroxide BaO2 burns, large amount of heat is
produced to initiate thermite process to produce molten iron.
 Thermite process is considered as a redox reaction.
 This process is highly exothermic, that is, it gives out a lot of heat during reaction, and it is used
for welding steel object such as railway lines.
Extraction of Tin from Its Ore
 Most important tin ore is cassiterite. Cassiterite contains SnO2 and unwanted materials such as
sand, soil, oil, sulphur and carbon.
 Concentration process
- Tin ore is concentrated by froth floatation method. Tin ore is crushed to fine powder and
mixed with water and special oils (frothing agents) in large tank.
- Mixture is agitated by blowing air to form froth. Unwanted materials sink to bottom of tank.
- Froth contains particles of concentrated tin ore and floats to top of tank where it is removed.
- Concentrated tin ore is then dried and roasted to remove impurities such as carbon, sulphur
and special oils.
 Reduction process
- Concentrated tin ore is mixed with coke.
- Mixture is heated to high temperature (about 1360 °C) in furnace
- During eating, tin(IV) oxide is reduced by carbon to molten tin and carbon is oxidised to
carbon dioxide and carbon monoxide.
SnO2 + C  Sn + CO2
SnO2 + 2C  Sn + 2CO
- Carbon monoxide produced can also reduce tin(IV) oxide.
SnO2 + 2CO  Sn + 2CO2
- Molten tin is then tapped off and poured into mould and solidified into ingots.

24. 24|여철우-화학
 Summary of extraction of tin
The Use of Carbon as the Main Reducing Agent in Metal Extraction
 Chemical reason: C is more reactive than Zn, Fe, Sn and Pb. Therefore, C can easily reduce oxides
of these metals
 Economical reason: C is cheap and can be obtained easily. Reduction of metal ores using coke is
cheaper than using electricity for electrolysis of molten ores.
 Environmental reason: CO2 produced during metal extraction is non-poisonous and does not
pollute atmosphere.
3.4 Redox Reactions in Electrolytic Cell and Chemical Cell
Electrolytic Cells
 Basic structure of electrolytic cell
- Battery (Supply electrical energy)
- Electrolyte (Supply free ions for conducting electric current)
- Two electrodes (Transfer of electrons)
 Electrolysis involves three main aspects
- External circuit
- Reaction occurring in electrolyte
- Reactions at electrodes
 When electrolysis occurs, electrical energy is converted into chemical energy.
 Electrical energy is converted into chemical energy. Electrical energy is used to decompose
electrolyte in electrolytic cell.

25. 여철우-화학|25
Chemical Cells
 Basic structure:
- Connecting wires (For electrons to flow through in external circuit)
- Electrolyte (For ions to flow through)
- Two electrodes (Transfer of electrons)
 Chemical energy is converted into electrical energy

26. 26|여철우-화학
Redox Reactions in Electrolytic Cells
 Ionic compounds in molten state dissolved in water are electrolytes.
 Oxidation occurs at anode, reduction occurs at cathode.
 In electrolytic cell, electrons flow from anode (positive electrode) to cathode (negative electrode)
through connecting wire.
Electrolysis of Molten Lead(II) Bromide
 When molten lead(II) bromide is electrolysed, cations are attracted to cathode and anions are
attracted to anode.
 Redox reaction occurs:
At cathode At anode
Pb2+ ions gain electrons to form lead metal
Pb2+ + 2e-  Pb (Reduction)
Br- ions lose electrons to form bromine
molecules 2Br-  Br2 + 2e- (Oxidation)
Overall reaction: PbBr2  Pb + Br2
Electrolysis of Copper(II) Sulphate Solution using Inert Electrodes
 Electrolysis of copper(II) sulphate uses Platinum electrodes as inert electrodes.
 Aqueous solution of copper(II) sulphate contains four types of ions
- CuSO4: Cu2+ and SO4
2-
- H2O: H+ and OH-
 Cu2+ and H+ ions are attracted to cathode and SO4
2- and OH_ ions are attracted to anode
 Redox reaction occurs:
At cathode At anode
Cu is below hydrogen in electrochemical series.
Hence Cu2+ are discharged are cathode.
Cu2+ ions are reduced to copper metal.
Cu2+ + 2e-  Cu
OH- is below SO4
2- ion in electrochemical
series. Hence, OH- ions are discharged at
anode.
OH- ions are oxidised to oxygen gas
4OH-  O2 + 2H2O + 4e-
Overall reaction: 2CuSO4 + 2H2O  2Cu + O2 + 2H2SO4

27. 여철우-화학|27
Electrolysis of Copper(II) Sulphate Solution using Copper Electrodes
 If electrolysis of copper(II) sulphate solution is carried out using reactive electrodes such as copper
electrodes, both OH- ions and SO4
2- ions are not discharged. Instead, copper anode dissolves to
form copper(II) ions.
 Redox reaction occurs:
At cathode At anode
Cu2+ + 2e- -> Cu (Reduction) Cu  Cu2+ + 2e- (Oxidation)
Overall reaction:
 Transfer of copper from anode to cathode.
 Concentration of CuSO4 change and blue colour of electrolyte does not fade.
Electrolysis of Concentrated Sodium Chloride Solution
 Aqueous solution of sodium chloride contains four types of ions
- NaCl: Na+ and Cl-
- H2O: H+ and OH-
 Na+ and H+ ions are attracted to cathode and Cl- and OH- ions are attracted to anode.
 Redox reaction occurs:
At cathode At anode
Hydrogen ions (H+) are discharged at cathode.
2H+ + 2e-  H2
H+ are reduced to hydrogen gas
Na+ ions remain in solution.
Chloride ions (Cl-) are discharged at anode.
2Cl-  Cl2 + 2e-
Cl- are oxidised to chlorine gas.
OH- ions remains in solution
Overall reaction: 2NaCl + 2H2O  2NaOH + H2 + Cl2
 Electrolysis of concentrated NaCl solution produces one volume of hydrogen at cathode, one
volume of chlorine at anode and sodium hydroxide solution.

28. 28|여철우-화학
Electrolysis of Dilute Sodium Chloride Solution
 Dilute sodium chloride solution contains:
- NaCl: Na+ and Cl-
- H2O: H+ and OH-
 Redox reaction occurs:
At cathode At anode
H+ ions gains electrons from cathode to form
hydrogen gas.
2H+ + 2e-  H2
OH- ions donate electrons to anode to form
oxygen gas and water.
4OH-  O2 + 2H2O + 2e-
Overall reaction: 2H2O  2H2 + O2
 Electrolysis of dilute NaCl solution produces two volumes of hydrogen at cathode and one
volume of oxygen at anode.
 Since water is being removed by decomposition to form H2 and O2, concentration of NaCl
increases gradually.
Redox Reactions in Daniell Cell
 Daniell cell is made up of zinc plate dipped into zinc sulphate solution and copper plate dipped
into copper(II) sulphate solution
 Function of porous pot / salt bridge is to
- Separate zinc sulphate solution from copper(II) sulphate solution so that solutions do not mix.
- Complete the electric circuit by allowing ions to pass through it.
 Zinc is more electropositive than copper, hence, zinc plate acts as negative electrode and copper
plate acts as positive electrode.
 At negative electrode (Zinc plate)
- Zinc is oxidised to zinc ions (Zn  Zn2+ + 2e-)
- Zinc metal acts as reducing g agent
 At positive electrode (Copper plate)
- Copper(II) ions gain electrons form zinc and is reduced to copper metal
- Copper(II) ions act as oxidising agent

29. 여철우-화학|29
 Overall reaction
- Redox reaction
Zn + Cu2+  Zn2+ + Cu
- Redox reactions that occur in Daniell cell and many other chemical cells are displacement
reactions.
- When Daniell cell is in use,
 Concentration of Zn2+ ions in solution increases
 Blue colour of CuSO4 solution fades gradually as more copper is deposited and
concentration of Cu2+ decreases.
 Mass of zinc electrode decrease gradually
 Mass of copper electrode increases gradually.
 Voltage of Daniell cell
- If concentrations of both ZnSO4 and CuSO4 solutions are 1.0 mol dm-3, maximum voltage of
Daniell cell is 1.10 V.
- Voltage of cell will decrease with time when cell is being used as concentration of Cu2+ ions
decreases.
Redox Reactions in Dry Cell
 Dry cell is made up of zinc container as anode and carbon rod as cathode. Electrolyte in dry cell
is paste consisting of ammonium chloride, zinc chloride and little water.
 When dry cell is used to generate electrical energy, oxidation occurs at negative terminal (zinc
container) and reduction occurs as negative terminal carbon rod).
 Redox reaction occurs:
At anode (negative terminal) At cathode (positive terminal)
Zn is oxidised to Zn2+.
Zn  Zn2+ + 2e-
Electrons flow from zinc container to carbon rod
NH4
+ is reduced to NH3 and H2
2NH4
+ + 2e-  2NH3 + H2
H2 gas produced is remove by reaction with
MnO2
2MnO2 + H2  Mn2O3 + H2O
Overall reaction: Zn + 2NH4
+ + 2MnO2  Zn2+ + 2NH3 + Mn2O3 + H2O
 Oxidising agent: Ammonium ion, NH4
+
 Reducing agent: Zinc

30. 30|여철우-화학
Redox Reactions in Alkaline Cell
 Alkaline cell is also known as alkaline battery.
 Negative terminal: Zinc container
 Positive terminal: Manganese(IV) oxide powder
 Electrolyte: LiOH or KOH
 Redox reaction occurs:
At anode (negative terminal) At cathode (positive terminal)
Zn is oxidised to Zn2+.
Zn  Zn2+ + 2e-
Electrons flow from zinc container to MnO2
MnO2 is reduced to Mn2O3
2MnO2 + H2O  Mn2O3 + 2OH-
Overall reaction: Zn + 2MnO4 + H2O  Zn2+ + Mn2O3 + 2OH-
 Oxidising agent: MnO2
 Reducing agent: Zinc
Redox Reactions in Mercury Cell
 Positive terminal, Mercury(II) oxide
 Negative metal: Zinc
 Electrolyte: KOH
 Redox reaction occurs
At anode (negative terminal) At cathode (positive terminal)
Zn is oxidised to Zn2+.
Zn + 2OH-  Zn(OH)2 + 2e-
Electrons flow from zinc electrode to HgO
HgO is reduced to mercury
HO + H2O + 2e-  Hg + 2OH-
Overall reaction: Zn + HgO + H2O  Zn(OH)2 + Hg
 Oxidising agent: Mercury(II) oxide
 Reducing agent: Zinc

31. 여철우-화학|31
Redox Reactions in Lead-acid Accumulator
 Lead-acid accumulator is often known as car battery.
 It is a battery chemical cell that can be recharged by passing current through it from external d.c.
supply.
 Positive terminal: Lead plate coated with PbO2
 Negative terminal: Lead plate
 Electrolyte: Sulphuric acid
 Redox reaction occurs:
At anode (negative terminal) At cathode (positive terminal)
(a) Pb is oxidised to Pb2+ ions with release of
electrons.
Pb  Pb2+ + 2e-
(b) Electrons given out at cathode flow through
external circuit to positive terminal.
(c) White precipitate is produced when Pb2+
ions react with SO4
2- ions in sulphuric acid to
form lead(II) sulphate.
Pb2+ + SO4
2-  PbSO4
(d) Negative electrode becomes white because
white solid lead(II) sulphate is deposited on
its surface.
(e) Overall reaction at negative electrode:
Pb  Pb2+ + 2e-
Pb2+ + SO4
2-  PbSO4
Pb + SO4
2-  PbSO4 + 2e-
(a) Lead(IV) oxide is reduced to Pb2+ ions by
accepting electrons.
PbO2 + 4H+ + 2e-  Pb2+ + 2H2O
(b) White solid is produced when Pb2+ ions
react with SO4
2- ions in sulphuric acid to
form lead(II) sulphate.
Pb2+ + SO4
2-  PbSO4
(c) White solid, lead(II) sulphate, then deposits
on surface of positive electrode to form
white coating.
(d) Overall reaction at positive electrode:
PbO2 + 4H+ + 2e-  Pb2+ +2H2O
Pb2+ + SO4
2-  PbSO4
PbO2 + 4H+ + SO4
2- + 2e-  PbSO4 +
2H2O
Overall reaction: Pb + PbO2 + 4H+ + 2SO4
2-  2PbSO4 + 2H2O
 Oxidising agent: Lead(IV) oxide
 Reducing agent: Lead
 During discharge, sulphuric acid is used up

32. 32|여철우-화학
Recharging of Lead-acid Accumulator
 Direct current is passed through it, in direction which is opposite to discharge.
 At negative terminal, lead(II) sulphate is reduced to lead
PbSO4 + 2e-  Pb + SO4
2-
 At positive terminal, lead(II) sulphate is oxidised to lead(IV) oxide
PbSO4 + 2H2O  PbO2 + 4H+ + SO4
2-
 Overall reaction during recharging:
2PbSO4 + 2H2O  Pb + PbO2 + 4H+ + 2SO4
2-
Compare and Contrast Electrolytic Cells and Chemical Cells in Terms of Redox Reactions
 In electrolytic cell, electrode connected to positive terminal of chemical cell is called anode.
Conversely, electrode connected to negative terminal of chemical cell is called anode.
 In both cells, oxidation occurs at anode, reduction occurs at cathode
 In electrolytic cells, anions from electrolyte donate electrons and is oxidised at anode.
 In chemical cell, more electropositive metal is oxidised at anode and donate electrons to cathode.
 Comparison of chemical and electrolytic cell in terms of redox reactions
Chemical cell (E.g.: Daniell cell)
Electrolytic cell (E.g.: Electrolysis of molten
NaCl)
At anode: Oxidation occurs
Zn  Zn2+ + 2e-
At anode: Oxidation occurs
2Cl- + Cl2 + 2e-
At cathode: Reduction occurs
Cu2+ + 2e-  Cu
At cathode: Reduction occurs
Na+ + e-  Na
 However, electrodes in chemical and electrolytic cells have different signs For example, anode in
chemical cell is negative electrode whereas anode in electrolytic cell is positive electrode.
Electrode Chemical cell Electrolytic cell
Anode
Negative terminal
Electrons are released at anode
Positive terminal
Electrons flow out from anode to battery.
Cathode
Positive terminal
Electrons are removed from cathode by
positive ions present in electrolyte
Negative terminal
Electrons flow from battery and enter
cathode

33. 여철우-화학|33
3.5 Appreciating the Ability of the Elements to Change Their Oxidation Numbers
Various Applications of the Changes of Oxidation Numbers in Substance
 Changes in oxidation number of substance can be applied in following processes:
- Extracting metal from its ore
- Corrosion of metal
- Preventing corrosion of metal
- Generation of electricity by cells
- Recycling of metals
 In extraction of iron from its ores, following changes in oxidation number of both iron and carbon
occur:
Fe2O3 + 3C  2Fe + 3CO
 In corrosion of iron, following changes in oxidation numbers occur:
4Fe + 3O2 + 2xH2O  2Fe2O3.xH2O
 Following chemical changes occur when zinc is used in prevention of rusting
Zn  Zn2+ + 2e-
O2 + 2H2O + 4e-  4OH-
Oxidation number of zinc changes from 0 to +2 while oxidation number of oxygen changes from
0 to –2.
 When Daniell cell is used to generate electricity, overall cell reaction is:
Zn + Cu2+  Zn2+ + Cu
Oxidation number of zinc changes from 0 to +2 while oxidation number of copper changes from
+2 to 0.
The Occurrence of Various Ores in Our Country
 Gold
- Gold mines are found in Pahang, Terengganu, Sabah and Sarawak
 Iron, bauxite (Al2O3) and ilmenite (FeTiO3)
- Iron mines are found in Johor, Kedah, Pahang and Perak
- Bauxite is aluminium ore and is found in Johor
- Ilmenite is a titanium ore and is found in Terengganu and Pulau Pinang
 Tin
- Tin ores are found in Perak and Selangor
 Coal
- Coal mines are found in Sarawak
 Kaolin and barite (BaSO4)
- Kaolin is a type of clay used for making ceramics. It is found in Johor and Perak
- Barite is the chief ore of barium. It is found in Kelantan and Terengganu.

34. 34|여철우-화학
The Contribution of the Metal Extraction Industry in Enhancing the National Economy
 Metal extraction industry provides job opportunities and lowers unemployment rate, Export
revenue from tin and other metals enhances national economy.
 Our country exports tin ingots to countries like JPN, USA and GBR. This will earn foreign exchange
which can be used for national development.
 Metal extraction industry produces metals as raw materials for many other industries, such as
motor and construction industries.
Methods for Rust Prevention
Method Objects
Paint
Big objects like motor vehicles, ships and steel
bridges
Oil and grease Tools and machine parts
Phosphoric acid (H3PO4) Bottom (Chassis of cars)
Galvanising (Zinc-plating) Buckets, „zinc‟ roof
Tin-plating Food cans
Chrome-plating Taps, bicycle handle bars, car bumpers
Block of magnesium of zinc (Sacrificial metals) Underground pipes, ships
Stainless steel Cutlery, surgical instruments
Chemical Cell as a Source of Electrical Energy
 Chemical cells are alternative sources of renewable energy.
 Fuel cell is device in which fuel is oxidised in chemical cell so as to produce electricity directly.
 In hydrogen-oxygen fuel cell, chemical energy from redox reaction between hydrogen and
oxygen to form water is used to generate electric current.
 Fuel cell are used to power electric cars, but they differ from usual chemical cells in two ways:
- Fuel and oxygen are fed into cell continuously
- Electrodes are made from inert material such as platinum that does not react during process.