Curriculum
Chemistry
Chapter 1 - Transition Elements and Iron Metallurgy
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1.1Where they are in the periodic table1.2The main transition elements (d-block)1.3The four transition series1.4Economic importance of the first-series elements1.5Electron configurations of the first series1.6Rules for filling1.7Configurations of the common ions1.8Why Fe(II) is easily oxidized to Fe(III) but Mn(II) is not1.9Oxidation-state rules for the series1.10Calculating oxidation states in compounds1.11Definition of a transition element
Chapter 2 - Chemical Analysis
0/221.1Revision of concepts and laws needed in this chapter1.2Importance1.3Kinds of analysis1.4Qualitative analysis1.5A. Identification of acid radicals (anions)1.6(1) Dilute hydrochloric acid group1.7(2) Concentrated sulphuric acid group1.8(3) Barium chloride group1.9Practical differentiation between pairs of salts (typical exam type)1.10Reasoning questions1.11Detailed Comparison Matrix: Diagnostic Silver Nitrate (AgNO3) Confirmatory Matrix
Chapter 3 - Chemical Equilibrium
0/221.1The equilibrium system1.2Types of chemical reactions1.3Chemical equilibrium1.4Rate of reaction1.5Factors affecting the rate of a reaction1.6Nature of the reactants1.7Concentration and the law of mass action1.8Temperature — collision theory1.9Pressure1.10Le Chatelier's principle1.11Catalysts1.12Light
2.1Strong and weak electrolytes2.2Hydronium ion2.3Definitions2.4Ostwald's dilution law2.5Hydronium ion concentration of a weak acid2.6Hydroxide ion concentration of a weak base2.7Ionization of water and the ionic product Kw2.8pH and pOH2.9Solubility and solubility product2.10Summary of key relations
Chapter 4 - Electrochemistry
0/331.1The starting experiment: zinc in copper sulphate1.2The idea behind the galvanic cell1.3Construction of the Daniell cell1.4What happens at each electrode1.5When does the current stop?1.6The salt bridge — what it is and why it is needed1.7Why we need a reference1.8The standard hydrogen electrode1.9Standard electrode potentials1.10Measuring an electrode potential (idea of the experiment)1.11Definition1.12What the series tells us (the conclusions to know)1.13Calculating the cell emf1.14Cell notation (cell diagram)
3.1The electrolytic cell3.2Example: electrolysis of copper(II) chloride solution3.3First law3.4Second law3.5The faraday3.6The working formula3.7Worked examples3.8Electroplating3.9Extraction of aluminium3.10Electrorefining of copper3.11Galvanic cells versus electrolytic cells3.12Scientists and the value of electrochemistry
Chapter 5 - Organic Chemistry
0/721.1Historical background1.2Modern definition1.3Why are there so many organic compounds?1.4Comparison: organic and inorganic compounds1.5Molecular formula1.6Structural formula1.7Isomerism1.8Detecting carbon and hydrogen in an organic compound1.9Classification of hydrocarbons1.10Structure and the homologous series1.11Alkyl radicals (R–)1.12IUPAC nomenclature of alkanes1.13Methane — occurrence and preparation1.14Physical properties1.15Chemical properties of alkanes
5.1Functional groups5.2Alcohols and phenols — the difference5.3Nomenclature of alcohols5.4Classification of alcohols5.5Ethanol — preparation5.6Denatured (methylated) alcohol — "red spirit"5.7Physical properties5.8Chemical properties of ethanol5.9Uses of alcohols5.10Phenols5.11Preparation5.12Physical properties5.13Chemical properties5.14Tests for phenol
6.1The carboxyl group6.2Types6.3Nomenclature6.4Preparation of acetic acid6.5Physical properties6.6Chemical properties6.7Aromatic acids: benzoic acid6.8Important acids and their uses6.9Definition and formation6.10Nomenclature6.11Preparation6.12Physical properties6.13Chemical properties6.14Applications of esters6.15Aspirin6.16Types of organic reactions (summary)6.17Revision key points for Part B
7.1The birth of organic chemistry and the vital-force theory7.2Why organic compounds are so abundant7.3Organic versus inorganic compounds7.4Molecular formula, structural formula and isomers7.5Detecting carbon and hydrogen in an organic compound7.6Hydrocarbons and the alkane series7.7The alkyl radical and IUPAC naming7.8Methane, physical and chemical properties of alkanes7.9Practice with answers7.10Ministry exam traps
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Text lesson
Lesson 2: Commercial Batteries, Fuel Cells and Metal Corrosion
Lesson 2: Commercial Batteries, Fuel Cells and Metal Corrosion
Overview
Overview
🌳 Concept tree
Lesson 2→Commercial Batteries, Fuel Cells and Metal Corrosion
├── Kinds of galvanic cells
├── Primary cells
- ├── Secondary cells (storage cells
- accumulators)
├── Why corrosion matters
├── Mechanism→corrosion is a galvanic cell
├── Factors that cause corrosion
└── Protection of iron
Lesson 2: Commercial Batteries, Fuel Cells and Metal Corrosion · 1 / 9
Kinds of galvanic cells
Kinds of galvanic cells
- According to whether the cell can be recharged, galvanic cells are classified as primary cells and secondary cells.
- (The fuel cell is a special case that is supplied continuously with its reactants.)
Lesson 2: Commercial Batteries, Fuel Cells and Metal Corrosion · 2 / 9
Primary cells
Primary cells
- A primary cell stores energy as chemical energy and releases it through a spontaneous, irreversible redox reaction.
- It stops working when the anode material is used up or the ions of the cathode half-cell are consumed.
- It cannot be recharged, because its reaction cannot be reversed in a practical way.
- For convenience, especially in portable devices, primary cells are made in the "dry" form, which gives a steady potential for a long time and allows small sizes.
- (A) The mercury cell. Made as a small cylinder or button, so it is used in hearing aids, watches and cameras.
| Comparison aspect | Main idea | Keywords |
|---|---|---|
| Negative pole (anode) | zinc | zinc |
| Positive pole (cathode) |
| mercuryoxidegraphite |
| Electrolyte |
| potassiumhydroxidetightlysealed |
| Cell reaction | Zn + HgO → ZnO + Hg | |
| • |
| practicallyconstantuntilcell |
| • |
| Mercurypoisonoususedcells |
| (B) The fuel cell. | Hydrogen burns in oxygen: 2H₂ + O₂ → 2H₂O + energy. In a fuel cell this reaction is carried out in a controlled way so that the energy appears directly as electricity instead of heat and light | Hydrogenburnsoxygen2H₂O |
| Comparison aspect | Main idea | Keywords |
|---|---|---|
| • |
| porouscarbonelectrodeshollow |
| Anode (oxidation) | 2H₂(g) + 4OH−(aq) → 4H₂O(l) + 4e− (E°ox = +0.83 V) | 4H₂O |
| Cathode (reduction) | O₂(g) + 2H₂O(l) + 4e− → 4OH−(aq) (E°red = +0.40 V) | 2H₂O |
| Total | 2H₂(g) + O₂(g) → 2H₂O emf = 1.23 V | 2H₂O |
- Features it is not consumed like other cells because fuel is supplied from outside and the product (water) is removed continuously; it does not store energy — it needs continuous supply of reactants.
- It works at high temperature, so the water leaves as vapour, which can be condensed and used as drinking water by astronauts.
- It is preferred in spacecraft because its fuel (hydrogen and oxygen) is the same fuel used to launch rockets, and it is light and clean.
Lesson 2: Commercial Batteries, Fuel Cells and Metal Corrosion · 3 / 9
Secondary cells (storage cells, accumulators)
Secondary cells (storage cells, accumulators)
- In a secondary cell the chemical reactions are reversible.
- During use (discharge) the cell works as a galvanic cell; it is recharged by passing a direct current from an outside source through it in the direction opposite to the discharge current, so it then works as an electrolytic cell and the chemical energy is stored again.
- (A) The lead–acid battery (car battery).
- A car battery has six cells in series, each giving about 2 V, so total emf = 6 × 2 = 12 V.
- Anode (negative plate) lead grid filled with spongy lead, Pb.
- Cathode (positive plate) lead grid filled with a paste of lead(IV) oxide, PbO₂.
- Plates are separated by insulating sheets and dip in dilute sulphuric acid (the electrolyte), all in a case of hard rubber or plastic that the acid does not attack.
- Discharge (galvanic cell):
AnodePb(s) + SO₄2−(aq) → PbSO₄(s) + 2e− (E°ox = +0.36 V)
CathodePbO₂(s) + 4H+(aq) + SO₄2−(aq) + 2e− → PbSO₄(s) + 2H₂O(l) (E°red = +1.69 V)
TotalPb + PbO₂ + 4H+ + 2SO₄2− → 2PbSO₄ + 2H₂O emf ≈ 2.05 V per cell.
- During discharge both plates are slowly covered with lead(II) sulphate, sulphuric acid is used up and water is produced, so the acid becomes more dilute and its density falls.
- The state of charge is checked with a hydrometer for a fully charged battery the acid density is about 1.28–1.30 g/cm3; if it falls below about 1.20 g/cm3 the battery needs recharging.
- Charging (electrolytic cell) a d.c.
- source of slightly higher voltage than the battery’s is connected so that its negative pole goes to the battery’s negative plate and its positive pole to the positive plate.
- The discharge reaction is reversed:
- 2PbSO₄(s) + 2H₂O(l) → Pb(s) + PbO₂(s) + 4H+(aq) + 2SO₄2−(aq)
- Lead sulphate is turned back into lead at one plate and lead(IV) oxide at the other, and the acid concentration is restored.
- In a car the dynamo/alternator recharges the battery continuously while the engine runs.
- (B) The lithium-ion battery.
- A dry, rechargeable battery used in mobile phones, laptops and modern (electric) cars in place of the lead–acid battery, because it is lighter, stores more energy per unit mass, and is small in size.
- It is based on lithium, which is the lightest metal and has a very low reduction potential (about −3.04 V), so it can give a high cell voltage.
- Inside the metal case three thin layers are rolled in a spiral: the anode (negative): lithium–graphite, LiC₆; the cathode (positive): cobalt(IV) oxide, CoO₂, which becomes lithium cobalt oxide, LiCoO₂, as the cell discharges; and a very thin plastic separator that keeps the electrodes apart but lets Li+ ions pass.
- The layers are soaked in an anhydrous (water-free) electrolyte, a solution of lithium hexafluorophosphate, LiPF₆, in an organic solvent.
- Discharge at the anode LiC₆ → C₆ + Li+ + e−; at the cathode CoO₂ + Li+ + e− → LiCoO₂.
- Total LiC₆ + CoO₂ → C₆ + LiCoO₂ emf ≈ 3 V.
- On recharging, the reactions run in the opposite direction and lithium ions move back through the separator.
Lesson 2: Commercial Batteries, Fuel Cells and Metal Corrosion · 4 / 9
Why corrosion matters
Why corrosion matters
- Corrosion is the chemical (electrochemical) attack on a metal by its environment.
- It causes huge economic losses; it is estimated that about a quarter of the world’s yearly iron production is lost to corrosion.
- The corrosion of iron is called rusting.
- This is why preventing it is considered a basis of protecting the world economy.
Lesson 2: Commercial Batteries, Fuel Cells and Metal Corrosion · 5 / 9
Mechanism: corrosion is a galvanic cell
Mechanism: corrosion is a galvanic cell
- Very pure iron does not rust easily.
- Industrial iron and steel contain impurities (for example carbon) and are alloys that are never perfectly uniform.
- Wherever a more active region (iron) touches a less active region (carbon or another metal) in the presence of moisture, a tiny local galvanic cell forms.
- The more active metal is the anode and is corroded; the less active one is the cathode.
- For iron in damp air:
| Comparison aspect | Main idea | Keywords |
|---|---|---|
| Anode |
| ironitself2Fe2passes |
| Cathode | (carbon/impurity regions): O₂(g) + 2H₂O(l) + 4e− → 4OH−(aq) | carbonimpurityregions2H₂O |
| Ions meet | 2Fe2+ + 4OH− → 2Fe(OH)₂(s) (iron(II) hydroxide) | 2Fe2ironhydroxide |
| Further oxidation by dissolved oxygen |
| ironhydroxideafterlosing |
| Overall | 2Fe + 3/2 O₂ + 3H₂O → 2Fe(OH)₃ | 3H₂O |
| • |
| Rustingslowpurewater |
Lesson 2: Commercial Batteries, Fuel Cells and Metal Corrosion · 6 / 9
Factors that cause corrosion
Factors that cause corrosion
| Comparison aspect | Main idea | Keywords |
|---|---|---|
| Factors in the metal itself |
| Heterogeneityalloysindustrialrarely |
| Factors in the surrounding medium |
| watermoistureoxygendissolved |
Lesson 2: Commercial Batteries, Fuel Cells and Metal Corrosion · 7 / 9
Protection of iron
Protection of iron
- The aim is usually to cover the iron so that it is isolated from air and water, or to make it the cathode of a cell.
| Comparison aspect | Main idea | Example | Keywords |
|---|---|---|---|
| Painting or oiling |
| — | coatinggreasevarnishsimple |
| Covering with a metal | Less active metal ( |
| Lessactiveplatingfood |
| Cathodic protection with a sacrificial anode | buried iron pipes and ships’ hulls (always in contact with salty water or damp soil) are connected to blocks of a more active metal |
| buriedironpipesships |
Lesson 2: Commercial Batteries, Fuel Cells and Metal Corrosion · 8 / 9
Part 9
Part 9
- 📖
🔍 In-Depth Explanations, Deductions & Multi-Book Insights:
Chemistry Thanaweya Amma (2026/2027)
- Lesson 2 – Commercial Batteries, Fuel Cells and Metal Corrosion
- Detailed Academic Lesson Overview:
- Batteries store chemical energy for on-demand electrical delivery.
- Primary cells (Mercury cell, Fuel cell) are irreversible.
- Secondary rechargeable cells (Lead-acid car battery, Lithium-ion battery) reverse their chemistry during charging.
- We also examine electrochemical metal corrosion and cathodic/anodic protection.
🌳 Hierarchy Classification & Conceptual Roadmap
Commercial Cells and Corrosion
├── 1. Primary Batteries (Non-rechargeable)
- │ ├── Mercury Cell (1.35 V): Zn anode, HgO cathode, KOH electrolyte; compact for hearing aids
- │ └── Hydrogen Fuel Cell (1.23 V): Porous carbon cylinders with H2 and O2; operates at high temp, provides drinking water in space
├── 2. Secondary Batteries (Rechargeable)
| Comparison aspect | Main idea | Keywords |
|---|---|---|
| • | │ ├── Lead-Acid Car Battery (12 V / 6 cells × 2 V) | LeadAcidBatterycells |
| │ │ ├── Discharging | Pb(s) + PbO2(s) + 2 H2SO4 → 2 PbSO4(s) + 2 H2O (acid density drops) | PbO2H2SO4PbSO4acid |
| │ │ └── Recharging | Connect to DC source with voltage slightly higher than 12 V | Connectsourcevoltageslightly |
| • | │ └── Lithium-Ion Battery (3.0 V) | LithiumBattery |
| │ ├── Anode |
| LithiumgraphiteLiC6Cathode |
| │ └── Electrolyte | Non-aqueous anhydrous LiPF6 (Lithium hexafluorophosphate) | aqueousanhydrousLiPF6Lithium |
└── 3. Metal Corrosion and Rusting of Iron
| Comparison aspect | Main idea | Example | Keywords |
|---|---|---|---|
| ├── Electrochemical mechanism |
| — | anodeimpuritiescarboncathode |
| ├── Overall rust reaction | 2 Fe + 3/2 O2 + 3 H2O → 2 Fe(OH)3 (Hydrated Iron(III) oxide) | — | HydratedIronoxide |
| • | └── Corrosion Protection | — | CorrosionProtection |
| ├── Galvanic / Anodic Protection | Coating iron with more active metal (Zinc / Sacrificial anode Mg) | — | Coatingironactivemetal |
| └── Cathodic Protection |
| — | Coatingironlessactive |
| • | 📋 Master Teacher's Blackboard Rules & Exam Secrets | — | MasterTeacherBlackboardRules |
| RULE 1 |
| — | FuelcellDOESstore |
| RULE 2 |
| — | Fullychargedleadacid |
| RULE 3 |
| — | Duringrechargingconnectbattery |
| EXAM TRAP | Anodic coating ( | galvanizing with Zn) protects iron even if scratched because zinc oxidizes preferentially. Cathodic coating (e.g. tin-plating with Sn) causes iron to rust FASTER if scratched because iron becomes the anode! | Anodiccoatinggalvanizingprotects |
| 📊 Detailed Comparison Matrix | Comparative Synthesis of Primary and Secondary Battery Cells | — | ComparativeSynthesisPrimarySecondary |
| Battery Type | Anode (-) | Cathode (+) | Electrolyte | Total Cell EMF |
|---|---|---|---|---|
| Mercury Cell | Zinc amalgam (Zn/Hg) | Mercuric oxide (HgO) | KOH (Alkaline) | 1.35 V |
| Hydrogen Fuel Cell | Porous carbon (H2 gas) | Porous carbon (O2 gas) | Hot KOH solution | 1.23 V |
| Lead-Acid Accumulator | Spongy Lead (Pb) | Lead dioxide (PbO2) | Dilute H2SO4 (1.28 g/cm3) | 12.0 V (6 × 2.0 V) |
| Lithium-Ion Battery | Lithium Graphite (LiC6) | Lithium Cobalt Oxide (LiCoO2) | LiPF6 in non-aqueous solvent | 3.00 V |
Lesson 2: Commercial Batteries, Fuel Cells and Metal Corrosion · 9 / 9
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