IGCSE Chemistry Index

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Summary: Complete index of IGCSE Chemistry (0620/0971) concepts organised by the official Cambridge syllabus. All content extracted and verified from real past papers, mark schemes, and the official syllabus. Tags: index igcse chemistry hub Created: 2026-07-14 Last Updated: 2026-07-23

📋 Official Syllabus: Cambridge IGCSE Chemistry 0620 Syllabus (2023-2025) — full subject content with clickable topic links 📝 Answers: Syllabus Answers — point-form study guide

Directory

1. States of Matter
2. Atoms, Elements and Compounds
3. Stoichiometry
4. Electrochemistry
5. Chemical Energetics
6. Chemical Reactions
7. Acids, Bases and Salts
8. The Periodic Table
9. Metals
10. Chemistry of the Environment
11. Organic Chemistry
12. Experimental Techniques and Chemical Analysis

1 States of Matter

1.1 Solids, Liquids and Gases

Core:

1. State the distinguishing properties of solids, liquids and gases
2. Describe the structures of solids, liquids and gases in terms of particle separation, arrangement and motion
3. Describe changes of state in terms of melting, boiling, evaporating, freezing and condensing
4. Describe the effects of temperature and pressure on the volume of a gas

Supplement:

5. Explain changes of state in terms of kinetic particle theory, including the interpretation of heating and cooling curves
6. Explain, in terms of kinetic particle theory, the effects of temperature and pressure on the volume of a gas

1.2 Diffusion

Core:

1. Describe and explain diffusion in terms of kinetic particle theory

Supplement:

2. Describe and explain the effect of relative molecular mass on the rate of diffusion of gases

Related pages: States of Matter — combined overview with past paper data

2 Atoms, Elements and Compounds

2.1 Elements, Compounds and Mixtures

Core:

1. Describe the differences between elements, compounds and mixtures — see Elements, Compounds and Mixtures

2.2 Atomic Structure and the Periodic Table

Core:

1. Describe the structure of the atom as a central nucleus containing neutrons and protons surrounded by electrons in shells
2. State the relative charges and relative masses of a proton, a neutron and an electron
3. Define atomic number as the number of protons in the nucleus of an atom
4. Define nucleon number as the total number of protons and neutrons in the nucleus of an atom
5. Determine the electronic configuration of elements and their ions with proton number 1 to 20
6. State that: (a) Group VIII noble gases have a full outer shell; (b) the number of outer shell electrons is equal to the group number in Groups I to VII; (c) the number of occupied electron shells is equal to the period number

2.3 Isotopes

Core:

1. Define isotopes as different atoms of the same element that have the same number of protons but different numbers of neutrons
2. Interpret and use symbols for atoms and ions

Supplement:

3. State that isotopes of the same element have the same chemical properties because they have the same electronic configuration
4. Calculate the relative atomic mass of an element from the relative masses and abundances of its isotopes

2.4 Ions and Ionic Bonds

Core:

1. Describe the formation of positive ions (cations) and negative ions (anions)
2. State that an ionic bond is a strong electrostatic attraction between oppositely charged ions
3. Describe the formation of ionic bonds between Group I and Group VII elements, including dot-and-cross diagrams
4. Describe the properties of ionic compounds (melting/boiling points, electrical conductivity)

Supplement:

5. Describe the giant lattice structure of ionic compounds
6. Describe the formation of ionic bonds between metallic and non-metallic elements
7. Explain the properties of ionic compounds in terms of structure and bonding

2.5 Simple Molecules and Covalent Bonds

Core:

1. State that a covalent bond is a shared pair of electrons leading to noble gas electronic configurations
2. Describe covalent bonding in simple molecules: H₂, Cl₂, H₂O, CH₄, NH₃ and HCl, using dot-and-cross diagrams
3. Describe the properties of simple molecular compounds

Supplement:

4. Describe covalent bonding in CH₃OH, C₂H₄, O₂, CO₂ and N₂
5. Explain the properties of simple molecular compounds in terms of weak intermolecular forces

2.6 Giant Covalent Structures

Core:

1. Describe the giant covalent structures of graphite and diamond
2. Relate the structures of graphite (lubricant, electrode) and diamond (cutting tools) to their uses

Supplement:

3. Describe the giant covalent structure of silicon(IV) oxide, SiO₂
4. Describe the similarity in properties between diamond and silicon(IV) oxide

2.7 Metallic Bonding

Supplement:

1. Describe metallic bonding as the electrostatic attraction between positive ions in a giant metallic lattice and a ‘sea’ of delocalised electrons
2. Explain the properties of metals: good electrical conductivity, malleability and ductility

Related pages (with past paper data): Atomic Structure — 64 Q+A pairs · Ions and Ionic Bonds — 27 Q+A pairs · Covalent Bonding — 19 Q+A pairs · Metallic Bonding — 20 Q+A pairs · Giant Structures — 18 Q+A pairs

3 Stoichiometry

3.1 Formulae

Core:

1. State the formulae of the elements and compounds named in the subject content
2. Define the molecular formula of a compound as the number and type of different atoms in one molecule
3. Deduce the formula of a simple compound from the relative numbers of atoms present in a model or a diagrammatic representation
4. Construct word equations and symbol equations to show how reactants form products, including state symbols

Supplement:

5. Define the empirical formula of a compound as the simplest whole number ratio of the different atoms or ions in a compound
6. Deduce the formula of an ionic compound from the relative numbers of the ions present in a model or a diagrammatic representation or from the charges on the ions
7. Construct symbol equations with state symbols, including ionic equations
8. Deduce the symbol equation with state symbols for a chemical reaction, given relevant information

3.2 Relative Masses of Atoms and Molecules

Core:

1. Describe relative atomic mass, Ar, as the average mass of the isotopes of an element compared to 1/12th of the mass of an atom of ¹²C
2. Define relative molecular mass, Mr, as the sum of the relative atomic masses. Relative formula mass, Mr, will be used for ionic compounds
3. Calculate reacting masses in simple proportions. Calculations will not involve the mole concept

3.3 The Mole and the Avogadro Constant

Core:

1. State that concentration can be measured in g / dm³ or mol / dm³

Supplement:

2. State that the mole, mol, is the unit of amount of substance and that one mole contains 6.02 × 10²³ particles, e.g. atoms, ions, molecules; this number is the Avogadro constant
3. Use the relationship amount of substance (mol) = mass (g) / molar mass (g / mol) to calculate: (a) amount of substance, (b) mass, (c) molar mass, (d) relative atomic mass or relative molecular/formula mass, (e) number of particles
4. Use the molar gas volume, taken as 24 dm³ at room temperature and pressure, r.t.p., in calculations involving gases
5. Calculate stoichiometric reacting masses, limiting reactants, volumes of gases at r.t.p., volumes of solutions and concentrations of solutions expressed in g / dm³ and mol / dm³, including conversion between cm³ and dm³
6. Use experimental data from a titration to calculate the moles of solute, or the concentration or volume of a solution
7. Calculate empirical formulae and molecular formulae, given appropriate data
8. Calculate percentage yield, percentage composition by mass and percentage purity, given appropriate data

Related pages (with past paper data): Relative Masses and Moles · Empirical and Molecular Formulae · Reacting Masses · Percentage Yield and Purity · Water of Crystallisation · Concentration and Gas Volumes · Chemical Equations and Calculations

4 Electrochemistry

4.1 Electrolysis

Core:

1. Define electrolysis as the decomposition of an ionic compound, when molten or in aqueous solution, by the passage of an electric current
2. Identify in simple electrolytic cells: (a) the anode as the positive electrode, (b) the cathode as the negative electrode, (c) the electrolyte as the molten or aqueous substance that undergoes electrolysis
3. Identify the products formed at the electrodes and describe the observations made during the electrolysis of: (a) molten lead(II) bromide, (b) concentrated aqueous sodium chloride, (c) dilute sulfuric acid using inert electrodes made of platinum or carbon/graphite
4. State that metals or hydrogen are formed at the cathode and that non-metals (other than hydrogen) are formed at the anode
5. Predict the identity of the products at each electrode for the electrolysis of a binary compound in the molten state
6. State that metal objects are electroplated to improve their appearance and resistance to corrosion
7. Describe how metals are electroplated

Supplement:

8. Describe the transfer of charge during electrolysis to include: (a) the movement of electrons in the external circuit, (b) the loss or gain of electrons at the electrodes, (c) the movement of ions in the electrolyte
9. Identify the products formed at the electrodes and describe the observations made during the electrolysis of aqueous copper(II) sulfate using inert carbon/graphite electrodes and when using copper electrodes
10. Predict the identity of the products at each electrode for the electrolysis of a halide compound in dilute or concentrated aqueous solution
11. Construct ionic half-equations for reactions at the anode (to show oxidation) and at the cathode (to show reduction)

4.2 Hydrogen–Oxygen Fuel Cells

Core:

1. State that a hydrogen–oxygen fuel cell uses hydrogen and oxygen to produce electricity with water as the only chemical product

Supplement:

2. Describe the advantages and disadvantages of using hydrogen–oxygen fuel cells in comparison with gasoline/petrol engines in vehicles

Related pages (with past paper data): Electrolysis · Electrolysis of Aqueous Solutions · Electrolysis Applications · Fuel Cells · Ionic Equations and Half Equations

5 Chemical Energetics

5.1 Exothermic and Endothermic Reactions

Core:

1. State that an exothermic reaction transfers thermal energy to the surroundings leading to an increase in the temperature of the surroundings
2. State that an endothermic reaction takes in thermal energy from the surroundings leading to a decrease in the temperature of the surroundings
3. Interpret reaction pathway diagrams showing exothermic and endothermic reactions

Supplement:

4. State that the transfer of thermal energy during a reaction is called the enthalpy change, ∆H, of the reaction. ∆H is negative for exothermic reactions and positive for endothermic reactions
5. Define activation energy, Ea, as the minimum energy that colliding particles must have to react
6. Draw and label reaction pathway diagrams for exothermic and endothermic reactions using information provided, to include: (a) reactants, (b) products, (c) enthalpy change of the reaction, ∆H, (d) activation energy, Ea
7. State that bond breaking is an endothermic process and bond making is an exothermic process and explain the enthalpy change of a reaction in terms of bond breaking and bond making
8. Calculate the enthalpy change of a reaction using bond energies

Related pages (with past paper data): Energetics of Reactions · Bond Energy Calculations

6 Chemical Reactions

6.1 Physical and Chemical Changes

Core:

1. Identify physical and chemical changes, and describe the differences between them

6.2 Rate of Reaction

Core:

1. Describe the effect on the rate of reaction of: (a) changing the concentration of solutions, (b) changing the pressure of gases, (c) changing the surface area of solids, (d) changing the temperature, (e) adding or removing a catalyst, including enzymes
2. State that a catalyst increases the rate of a reaction and is unchanged at the end of a reaction
3. Describe practical methods for investigating the rate of a reaction including change in mass of a reactant or a product and the formation of a gas
4. Interpret data, including graphs, from rate of reaction experiments

Supplement:

5. Describe collision theory in terms of: (a) number of particles per unit volume, (b) frequency of collisions between particles, (c) kinetic energy of particles, (d) activation energy, Ea
6. Describe and explain the effect on the rate of reaction of changing conditions using collision theory
7. State that a catalyst decreases the activation energy, Ea, of a reaction
8. Evaluate practical methods for investigating the rate of a reaction

6.3 Reversible Reactions and Equilibrium

Core:

1. State that some chemical reactions are reversible as shown by the symbol ⇌
2. Describe how changing the conditions can change the direction of a reversible reaction for: (a) the effect of heat on hydrated compounds, (b) the addition of water to anhydrous compounds limited to copper(II) sulfate and cobalt(II) chloride

Supplement:

3. State that a reversible reaction in a closed system is at equilibrium when: (a) the rate of the forward reaction is equal to the rate of the reverse reaction, (b) the concentrations of reactants and products are no longer changing
4. Predict and explain, for a reversible reaction, how the position of equilibrium is affected by: (a) changing temperature, (b) changing pressure, (c) changing concentration, (d) using a catalyst using information provided
5. State the symbol equation for the production of ammonia in the Haber process, N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
6. State the sources of the hydrogen (methane) and nitrogen (air) in the Haber process
7. State the typical conditions in the Haber process as 450 °C, 20 000 kPa / 200 atm and an iron catalyst
8. State the symbol equation for the conversion of sulfur dioxide to sulfur trioxide in the Contact process, 2SO₂(g) + O₂(g) ⇌ 2SO₃(g)
9. State the sources of the sulfur dioxide (burning sulfur or roasting sulfide ores) and oxygen (air) in the Contact process
10. State the typical conditions for the conversion of sulfur dioxide to sulfur trioxide in the Contact process as 450 °C, 200 kPa / 2 atm and a vanadium(V) oxide catalyst
11. Explain, in terms of rate of reaction and position of equilibrium, why the typical conditions stated are used in the Haber process and in the Contact process, including safety considerations and economics

6.4 Redox

Core:

1. Use a Roman numeral to indicate the oxidation number of an element in a compound
2. Define redox reactions as involving simultaneous oxidation and reduction
3. Define oxidation as gain of oxygen and reduction as loss of oxygen
4. Identify redox reactions as reactions involving gain and loss of oxygen
5. Identify oxidation and reduction in redox reactions

Supplement:

6. Define oxidation in terms of: (a) loss of electrons, (b) an increase in oxidation number
7. Define reduction in terms of: (a) gain of electrons, (b) a decrease in oxidation number
8. Identify redox reactions as reactions involving gain and loss of electrons
9. Identify redox reactions by changes in oxidation number using: (a) the oxidation number of elements in their uncombined state is zero, (b) the oxidation number of a monatomic ion is the same as the charge on the ion, (c) the sum of the oxidation numbers in a compound is zero, (d) the sum of the oxidation numbers in an ion is equal to the charge on the ion
10. Identify redox reactions by the colour changes involved when using acidified aqueous potassium manganate(VII) or aqueous potassium iodide
11. Define an oxidising agent as a substance that oxidises another substance and is itself reduced
12. Define a reducing agent as a substance that reduces another substance and is itself oxidised
13. Identify oxidising agents and reducing agents in redox reactions

Related pages (with past paper data): Rates of Reaction · Reversible Reactions · Haber Process and Ammonia · Redox Reactions

7 Acids, Bases and Salts

See full syllabus for complete details.

7.1 The characteristic properties of acids and bases

Core:

1. Describe the characteristic properties of acids in terms of their reactions with: (a) metals, (b) bases, (c) carbonates
2. Describe acids in terms of their effect on: (a) litmus, (b) thymolphthalein, (c) methyl orange
3. State that bases are oxides or hydroxides of metals and that alkalis are soluble bases
4. Describe the characteristic properties of bases in terms of their reactions with: (a) acids, (b) ammonium salts
5. Describe alkalis in terms of their effect on: (a) litmus, (b) thymolphthalein, (c) methyl orange
6. State that aqueous solutions of acids contain H+ ions and aqueous solutions of alkalis contain OH- ions
7. Describe how to compare hydrogen ion concentration, neutrality, relative acidity and relative alkalinity using universal indicator paper
8. Describe the neutralisation reaction: H+(aq) + OH-(aq) → H2O(l)

Supplement:

• Define acids as proton donors and bases as proton acceptors
• Define a strong acid and a weak acid in terms of dissociation
• State that hydrochloric acid is a strong acid: HCl(aq) → H+(aq) + Cl-(aq)
• State that ethanoic acid is a weak acid: CH3COOH(aq) ⇌ H+(aq) + CH3COO-(aq)

7.2 Oxides

1. Classify oxides as acidic (SO2, CO2) or basic (CuO, CaO)
2. Describe amphoteric oxides — react with both acids and bases
3. Classify Al2O3 and ZnO as amphoteric oxides

7.3 Preparation of salts

1. Describe preparation of soluble salts by reaction of an acid with: (a) alkali by titration, (b) excess metal, (c) excess insoluble base, (d) excess insoluble carbonate
2. Describe general solubility rules for salts
3. Define hydrated and anhydrous substances
4. Describe preparation of insoluble salts by precipitation
5. Define water of crystallisation (CuSO4•5H2O, CoCl2•6H2O)

Related pages: Acids and Bases — 126 Q+A · Indicators and pH — 90 Q+A · Making Salts · Salt Preparation Methods · Neutralization — 86 Q+A · Qualitative Analysis — 92 Q+A · Solubility Rules — 86 Q+A · Oxides · Amphoteric Oxide · Titration

8 The Periodic Table

See full syllabus for complete details.

8.1 Arrangement of elements

1. Describe the Periodic Table — elements arranged in periods and groups by increasing atomic number
2. Describe change from metallic to non-metallic character across a period
3. Describe relationship between group number and charge of ions formed
4. Explain similar chemical properties of same-group elements in terms of electronic configuration
5. Explain how position in the Periodic Table predicts properties
6. Identify trends in groups given information

8.2 Group I properties

1. Describe Group I alkali metals (Li, Na, K) — soft metals, trends: decreasing melting point, increasing density, increasing reactivity
2. Predict properties of other Group I elements

8.3 Group VII properties

1. Describe Group VII halogens (Cl, Br, I) as diatomic non-metals — trends: increasing density, decreasing reactivity
2. State appearance: Cl (pale yellow-green gas), Br (red-brown liquid), I (grey-black solid)
3. Describe displacement reactions of halogens
4. Predict properties of other Group VII elements

8.4 Transition elements

1. Describe transition elements: high densities, high melting points, form coloured compounds, act as catalysts
2. Describe variable oxidation numbers including Fe(II) and Fe(III)

8.5 Noble gases

1. Describe Group VIII noble gases as unreactive, monatomic gases — explained by electronic configuration

Related pages: Periodic Table Trends — 64 Q+A · Group 1 Alkali Metals — 38 Q+A · Group 7 Halogens — 11 Q+A · Group 0 Noble Gases — 14 Q+A · Transition Metals — 140 Q+A · Periodic Table · Metallic Character · Diatomic

9 Metals

See full syllabus for complete details.

9.1 Properties of metals

1. Compare physical properties of metals and non-metals: thermal conductivity, electrical conductivity, malleability and ductility, melting/boiling points
2. Describe chemical properties of metals: reactions with dilute acids, cold water/steam, and oxygen

9.2 Uses of metals

1. Describe uses of metals: Al in aircraft (low density), Al in cables (low density + good conductor), Al in food containers (corrosion resistant), Cu in wiring (good conductor + ductile)

9.3 Alloys and their properties

1. Describe alloys: Brass (Cu + Zn), Stainless Steel (Fe + Cr + Ni + C)
2. State alloys are harder and stronger than pure metals
3. Uses: stainless steel in cutlery (hardness + rust resistance)
4. Identify alloy structures from diagrams
5. Explain: different sized atoms → layers cannot slide → harder and stronger

9.4 Reactivity series

1. State reactivity series order: K > Na > Ca > Mg > Al > (C) > Zn > Fe > (H) > Cu > Ag > Au
2. Describe reactions with: (a) cold water, (b) steam, (c) dilute HCl — related to position in series
3. Deduce order of reactivity from experimental results
4. Relative reactivities in terms of tendency to form positive ions — displacement reactions
5. Explain aluminium’s apparent unreactivity — oxide layer

9.5 Corrosion of metals

1. State conditions for rusting: oxygen AND water → hydrated iron(III) oxide
2. Common barrier methods: painting, greasing, plastic coating
3. Describe how barrier methods prevent rusting — exclude O2 or H2O
4. Galvanising with zinc = barrier + sacrificial protection
5. Explain sacrificial protection in terms of reactivity series and electron loss

9.6 Extraction of metals

1. Ease of obtaining metals from ores related to position in reactivity series
2. Describe extraction of iron from hematite: burning coke, CO production, Fe2O3 reduction, Limestone decomposition, slag formation
3. Main ore of aluminium is bauxite — extracted by electrolysis
4. State equations for iron extraction
5. Aluminium extraction: role of cryolite, carbon anode replacement, electrode reactions

Related pages: Reactivity of Metals — 124 Q+A · Metal Extraction — 124 Q+A · Blast Furnace · Uses of Metals — 124 Q+A · Rusting of Iron — 125 Q+A · Alloy · Stainless Steel · Brass · Sacrificial Protection · Barrier Method

10 Chemistry of the Environment

See full syllabus for complete details.

10.1 Water

1. Chemical tests for water: anhydrous cobalt(II) chloride (blue → pink), anhydrous copper(II) sulfate (white → blue)
2. Test purity of water using melting point and boiling point
3. Distilled water used in practical chemistry — fewer impurities than tap water
4. Substances in natural water sources: dissolved O2, metal compounds, plastics, sewage, microbes, nitrates, phosphates
5. Some beneficial: dissolved O2 (aquatic life), some minerals essential
6. Some harmful: toxic metals, plastics, sewage microbes, nitrates/phosphates → deoxygenation
7. Domestic water treatment: sedimentation and filtration, carbon (tastes/odours), chlorination (kills microbes)

10.2 Fertilisers

1. Ammonium salts and nitrates used as fertilisers
2. NPK fertilisers provide nitrogen, phosphorus, potassium for plant growth

10.3 Air quality and climate

1. Clean air: ~78% N2, ~21% O2, remainder = noble gases + CO2
2. Sources of air pollutants: CO2 (complete combustion), CO + particulates (incomplete combustion), CH4 (decomposition/digestion), NOx (car engines), SO2 (fossil fuels containing S)
3. Adverse effects: CO2 → global warming; CO → toxic; particulates → respiratory problems; CH4 → global warming; NOx → acid rain, smog; SO2 → acid rain
4. Strategies: climate change (planting trees, less fossil fuels, renewable energy); acid rain (catalytic converters, low-S fuels, flue gas desulfurisation)
5. Photosynthesis: CO2 + H2O → glucose + O2 (chlorophyll, light)
6. How greenhouse gases cause warming: absorb/reflect/emit thermal energy, reduce heat loss to space
7. NOx formation in car engines; removal by catalytic converters: 2CO + 2NO → 2CO2 + N2

Related pages: Composition of Air — 39 Q+A · Air Pollution — 129 Q+A · Climate Change and Greenhouse Gases — 55 Q+A · Water Treatment — 40 Q+A · Sulfur and Compounds — 36 Q+A · Carbonates — 30 Q+A · Photosynthesis · NPK Fertiliser

11 Organic Chemistry

See full syllabus for complete details.

11.1 Formulae, functional groups and terminology

1. Draw/interpret displayed formulae — show all atoms and bonds
2. Write/interpret general formulae: alkanes CnH2n+2, alkenes CnH2n, alcohols CnH2n+1OH, carboxylic acids CnH2n+1COOH
3. Functional group = atom/group determining chemical properties of a homologous series
4. Homologous series = family of compounds with same functional group → similar chemical properties
5. Saturated = all C–C bonds are single; Unsaturated = one or more C=C bonds
6. Structural formula — unambiguous description of atom arrangement
7. Structural isomers — same molecular formula, different structural formula
8. Homologous series characteristics: same functional group, same general formula, differ by —CH2—, trend in physical properties, similar chemical properties

11.2 Naming organic compounds

1. Name/draw displayed formulae of: methane, ethane, ethene, ethanol, ethanoic acid, and reaction products
2. Identify compound type from name ending: -ane, -ene, -ol, -oic acid
3. Name/draw structural and displayed formulae of unbranched alkanes, alkenes, alcohols, carboxylic acids (up to C4)
4. Name/draw displayed formulae of unbranched esters (up to C4 alcohol + C4 acid)

11.3 Fuels

1–7. Fossil fuels, petroleum as hydrocarbon mixture, fractional distillation, fraction properties and uses

11.4 Alkanes

1. Bonding in alkanes: single covalent, saturated hydrocarbons
2. Properties: generally unreactive except combustion and substitution by chlorine
3. Substitution: one atom replaced by another; alkane + Cl2 (UV light) → chloroalkane + HCl

11.5 Alkenes

1. Bonding in alkenes: C=C double covalent bond, unsaturated hydrocarbons
2. Manufacture by cracking larger alkanes (high temp + catalyst)
3. Reasons for cracking: produce shorter alkanes (more useful as fuels) + alkenes (chemical feedstock)
4. Test for unsaturation: react with aqueous bromine → orange/brown to colourless
5. Addition reactions: only one product formed — (a) Br2, (b) H2 (Ni catalyst), (c) steam (H3PO4 catalyst, 300°C, 60 atm)

11.6 Alcohols

1. Manufacture of ethanol: (a) fermentation (glucose, yeast, 25–35°C, anaerobic) vs (b) hydration of ethene (300°C, 60 atm, H3PO4)
2. Combustion of ethanol — clean-burning fuel
3. Uses: solvent, fuel
4. Advantages/disadvantages of fermentation vs hydration

11.7 Carboxylic acids

1. Ethanoic acid reactions with metals, bases, carbonates — naming salts produced
2. Formation by oxidation of ethanol: (a) acidified KMnO4, (b) bacterial oxidation (vinegar)
3. Reaction with alcohol (acid catalyst) → ester + water

11.8 Polymers

1–13. Polymers = large molecules from monomers. Addition Polymerisation (no byproduct): poly(ethene). Condensation Polymerisation (H2O eliminated): Nylon (polyamide), PET (polyester). Proteins = natural polyamides. Thermoplastic vs Thermosetting. Environmental challenges and disposal.

Related pages: Organic Chemistry Fundamentals — 71 Q+A · Alkanes — 88 Q+A · Alkenes — 68 Q+A · Alcohols — 65 Q+A · Carboxylic Acids — 65 Q+A · Esters — 65 Q+A · Polymers — 66 Q+A · Fuels and Petroleum — 68 Q+A · Addition Polymerisation · Condensation Polymerisation · Nylon · PET · Proteins · Thermoplastic · Thermosetting

12 Experimental Techniques and Chemical Analysis

See full syllabus for complete details.

12.1 Experimental design

1. Name appropriate apparatus: stopwatches, thermometers, balances, burettes, pipettes, measuring cylinders, gas syringes
2. Advantages/disadvantages of experimental methods
3. Define: solvent, solute, solution, saturated solution, residue, filtrate

12.2 Acid-base titrations

1–2. Acid-base titration using burette, volumetric pipette, indicator; identify end-point

12.3 Chromatography

1–4. Paper Chromatography — separate coloured/colourless substances; locating agents; Rf values; pure vs impure substances

12.4 Separation and purification

1–3. Separation methods: solvent, filtration, crystallisation, simple distillation, fractional distillation; choose method given information; assess purity by melting/boiling point

12.5 Identification of ions and gases

1. Anion tests: CO32- (acid + limewater), Cl-/Br-/I- (HNO3 + AgNO3), NO3- (Al + NaOH → NH3), SO42- (HNO3 + Ba(NO3)2), SO32- (acidified KMnO4)
2. Cation tests with NaOH(aq) and NH3(aq): Al3+, NH4+, Ca2+, Cr3+, Cu2+, Fe2+, Fe3+, Zn2+
3. Gas tests: NH3 (damp red litmus → blue), CO2 (limewater → milky), Cl2 (damp litmus → red then bleaches), H2 (lighted splint → squeaky pop), O2 (glowing splint → relights), SO2 (acidified KMnO4 → colourless)
4. Flame tests: Li+ (red), Na+ (yellow), K+ (lilac), Ca2+ (orange-red), Ba2+ (apple green), Cu2+ (blue-green)

Related pages: Experimental Techniques · Separation Techniques · Paper Chromatography · Chemical Tests — 138 Q+A · Flame Test · Rf Value · Locating Agent · Purity · Titration

Sources

• BBC Bitesize GCSE Chemistry — IGCSE Chemistry topics, BBC (free educational resource)
• OpenStax Chemistry 2e — Chemistry 2e textbook, Rice University (free, CC BY 4.0)
• Cambridge IGCSE Chemistry 0620 — Syllabus 2023–2025, Cambridge Assessment International Education
• CK-12 Chemistry for High School — CK-12 Chemistry, CK-12 Foundation (free, CC BY-NC 3.0)
• ZNotes IGCSE Chemistry — Chemistry 0620 notes, ZNotes (free educational resource)

Data Sources

• 38 paired question papers + mark schemes (2015–2024)
• 1,229 Q+A pairs classified into 52 topics
• MS-driven extraction from chem-exam-data.js and chem-content-db.js

Related Indexes

• Chem-Index
• CS-Index