Curriculum
Chemistry
Chapter 1 - Transition Elements and Iron Metallurgy
0/28
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
More lessons
0/23
Text lesson
Lesson 4: Iron Chemical Properties, Reactions and Oxides
Lesson 4: Iron Chemical Properties, Reactions and Oxides
Overview
Overview
🌳 Concept tree
- Iron Chemical Properties, Reactions and Oxides
- ├── Physical properties
- ├── Chemical properties
- ├── Iron oxides
- ├── Summary of the oxides
- └── Typical "give reasons" for this chapter
Lesson 4: Iron Chemical Properties, Reactions and Oxides · 1 / 7
Physical properties
Physical properties
- Pure iron has no industrial importance because it is rather soft, malleable, ductile and magnetic; m.p. 1538 °C; density 7.87 g/cm3.
- Its properties depend on its purity and on the nature of the impurities.
- Many steels and iron alloys are made with specific properties.
Lesson 4: Iron Chemical Properties, Reactions and Oxides · 2 / 7
Chemical properties
Chemical properties
- Iron does not reach the maximum state (+8) that would correspond to removing all 8 electrons of 4s and 3d.
- States above +3 are unimportant; +2 comes from the 4s electrons and +3 gives the stable half-filled 3d5.
| Comparison aspect | Main idea | Keywords |
|---|---|---|
| Air | red-hot iron + dry air/oxygen → magnetic oxide: 3Fe + 2O₂ → Fe₃O₄. (In damp air at ordinary temperature iron rusts to hydrated iron(III) oxide.) | ironoxygenmagneticoxide |
| Water vapour | at red heat (about 500 °C): 3Fe + 4H₂O(steam) ⇌ Fe₃O₄ + 4H₂ | heat4H₂OsteamFe₃O₄ |
| Non-metals |
| 3Cl₂2FeCl₃heat |
| Acids |
| DiluteH₂SO₄giveiron |
| Comparison aspect | Main idea | Keywords |
|---|---|---|
| Worked example 7. |
| ironheatedsulphurMoles |
| Worked example 8. |
| Volumehydrogenobtainediron |
Lesson 4: Iron Chemical Properties, Reactions and Oxides · 3 / 7
Iron oxides
Iron oxides
1. Iron(II) oxide, FeO→(black powder, insoluble in water)
| Comparison aspect | Main idea | Keywords |
|---|---|---|
| Preparation |
| heatingironoxalatewithout |
| Properties | easily oxidized in hot air: 4FeO + O₂ → 2Fe₂O₃. Reacts with acids to give iron(II) salt and water: FeO + H₂SO₄ → FeSO₄ + H₂O | easilyoxidized4FeO2Fe₂O₃ |
| 2. Iron(III) oxide, Fe₂O₃ |
| brownfoundnaturehematite |
| Comparison aspect | Main idea | Keywords |
|---|---|---|
| Preparation |
| addingalkaliironsalt |
| Properties |
| insolublewaterusedpigment |
| 3. Black (magnetic) iron oxide, Fe₃O₄ | (found as magnetite) | foundmagnetite |
| Comparison aspect | Main idea | Keywords |
|---|---|---|
| • | A mixed oxide of iron(II) and iron(III) (FeO·Fe₂O₃) | mixedoxideironFe₂O₃ |
| Preparation |
| actionsteamironpartial |
| Properties |
| stronglymagneticconcentratedacids |
| Worked example 9 (the mixed oxide). |
| Fe₃O₄dissolvedexcessconcentrated |
Lesson 4: Iron Chemical Properties, Reactions and Oxides · 4 / 7
Summary of the oxides
Summary of the oxides
| Oxide | Colour | Occurs as | Fe oxidation state | With hot conc. acid gives |
|---|---|---|---|---|
| FeO | black | — | +2 | Fe(II) salt + water |
| Fe₂O₃ | red-brown | hematite | +3 | Fe(III) salt + water |
| Fe₃O₄ | black, magnetic | magnetite | +2 and +3 (mixed) | Fe(II) and Fe(III) salts |
Lesson 4: Iron Chemical Properties, Reactions and Oxides · 5 / 7
Typical "give reasons" for this chapter
Typical "give reasons" for this chapter
| Comparison aspect | Main idea | Keywords |
|---|---|---|
| Chromium and copper have abnormal electronic configurations | a half-filled (3d5) or full (3d10) sublevel makes the atom more stable | halffilledfull3d10 |
| Chromium is active but resists atmospheric attack | a non-porous oxide layer of larger volume than the metal forms on its surface | porousoxidelayerlarger |
| Cu, Ag and Au are transition elements |
| ionsincompletesublevels |
| • | The decrease in atomic radius across the first series is small: increasing nuclear attraction is opposed by increasing repulsion between 3d electrons | decreaseatomicradiusacross |
| Transition elements have high melting and boiling points | strong metallic bonds involving both 4s and 3d electrons | strongmetallicbondsinvolving |
| Most transition-element compounds are attracted by a magnet | they contain unpaired d electrons | containunpairedelectrons |
| • | Iron reacts with dilute acids to give iron(II) not iron(III) salts: the liberated hydrogen reduces Fe3+ to Fe2+ | Ironreactsdiluteacids |
| Scandium does not form Sc2+ | it loses all three outer electrons to reach the stable argon structure | losesthreeouterelectrons |
| Iron(III) chloride is paramagnetic | Fe3+ is 3d5 with five unpaired electrons | fiveunpairedelectrons |
| Concentrated nitric acid makes iron passive | a thin protective oxide layer forms on the surface | thinprotectiveoxidelayer |
| Most transition elements are catalysts |
| electronsformtemporarybonds |
Lesson 4: Iron Chemical Properties, Reactions and Oxides · 6 / 7
Part 7
Part 7
- 📖
🔍 In-Depth Explanations, Deductions & Multi-Book Insights:
Chemistry Thanaweya Amma (2026/2027)
- Lesson 4 – Iron Chemical Properties, Reactions and Oxides
- Detailed Academic Lesson Overview:
- Iron reacts with non-metals (S, Cl2), dilute acids to produce Iron(II) salts and hydrogen, and concentrated acids.
- We examine the preparation, oxidation, and acid behaviors of FeO, Fe2O3, and magnetic iron oxide Fe3O4.
🌳 Hierarchy Classification & Conceptual Roadmap
Iron and Iron Oxides Chemistry
├── 1. Pure Iron Reactions
- │ ├── 2 Fe + 3 Cl2 → 2 FeCl3 (Chlorine is a strong oxidizing agent, producing Fe3+)
│ ├── Fe + S → FeS (Sulfur produces Fe2+)
- │ ├── Fe + dil. H2SO4 → FeSO4 + H2 (Hydrogen is a reducing agent, preventing Fe3+ formation)
- │ └── Passivity Reaction with conc. HNO3 forms a microscopically thin, non-porous protective oxide layer
├── 2. Iron(II) Oxide (FeO)
- │ ├── Prepared by thermal decomposition of iron(II) oxalate in absence of air
- │ └── Reacts with dilute acids to give Fe(II) salts and water
├── 3. Iron(III) Oxide (Fe2O3 – Hematite)
- │ ├── Prepared by heating iron(II) sulfate (FeSO4) or iron(III) hydroxide (Fe(OH)3)
- │ └── Reacts only with hot concentrated acids to give Fe(III) salts
└── 4. Magnetic Iron Oxide (Fe3O4 – Magnetite)
| Comparison aspect | Main idea | Keywords |
|---|---|---|
| ├── Mixed oxide | FeO · Fe2O3 | Fe2O3 |
| • | └── Reacts with hot conc. H2SO4 to produce a mixture of Fe(II) and Fe(III) sulfates | ReactsconcH2SO4produce |
| • | 📋 Master Teacher's Blackboard Rules & Exam Secrets | MasterTeacherBlackboardRules |
| RULE 1 | Iron reacts with dilute acids to give Fe2+ salts because the liberated H2 gas is a reducing agent | Ironreactsdiluteacids |
| RULE 2 | Iron reacts with chlorine gas to give Fe3+ (FeCl3) because chlorine is a strong oxidizing agent | Ironreactschlorinegive |
| RULE 3 |
| Fe3O4mixedoxideproof |
| TRAP |
| Heatingironoxalateyields |
| 📊 Detailed Comparison Matrix | Diagnostic Reaction Matrix of Iron Oxides with Acids | DiagnosticReactionIronOxides |
| Oxide | Reaction with Dilute Acids | Reaction with Hot Conc. H2SO4 | Behavior on Heating in Air |
|---|---|---|---|
| Iron(II) oxide (FeO) | Dissolves giving Fe2+ salt + H2O | Reacts giving Fe2+ salt | Oxidizes rapidly to Fe2O3 (Hematite) |
| Iron(III) oxide (Fe2O3) | No reaction (Does not dissolve) | Dissolves giving Fe2(SO4)3 + H2O | Thermodynamically stable (No change) |
| Magnetite (Fe3O4) | No reaction (Does not dissolve) | Gives mixed FeSO4 + Fe2(SO4)3 + H2O | Oxidizes at high temp to Fe2O3 |
Lesson 4: Iron Chemical Properties, Reactions and Oxides · 7 / 7
Well done! You finished this lesson
What would you like to do next?