This is the complete, objective-by-objective curriculum guide for Cambridge IGCSE Chemistry (0620), covering every single Core and Supplement learning objective from the official syllabus. SECTION 1: STATES OF MATTER 1.1 Solids, liquids and gases - Distinguishing Properties: - Solids: Fixed shape and fixed volume; cannot flow; incompressible. - Liquids: No fixed shape (takes container shape); fixed volume; can flow; incompressible. - Gases: No fixed shape or volume (fills any container); flows easily; highly compressible. - Particle Structure & Arrangement: - Solid: Particles tightly packed in a regular 3D lattice, touching, vibrating about fixed positions. - Liquid: Particles mostly touching, random arrangement, slide past one another. - Gas: Particles far apart, random arrangement, moving rapidly and randomly in all directions. - Changes of State: - Melting: Solid \rightarrow Liquid (absorbs heat). - Freezing: Liquid \rightarrow Solid (releases heat). - Boiling: Liquid \rightarrow Gas at a specific, fixed boiling point throughout the liquid. - Evaporation: Liquid \rightarrow Gas occurring only at the liquid surface below the boiling point. - Condensation: Gas \rightarrow Liquid. - Kinetic Theory Explanation of State Changes: - Heating provides thermal energy \rightarrow converted to kinetic energy of particles \rightarrow particles vibrate/move faster. - At the melting/boiling point, temperature remains constant because energy is used to overcome/break intermolecular attractive forces, not to raise temperature. - Heating and Cooling Curves: - Slanted lines: Single state being heated/cooled; temperature changes as kinetic energy increases/decreases. - Horizontal plateaus: Phase change in progress; temperature remains constant as latent heat is absorbed or released to break or form intermolecular forces. - Gas Volume, Temperature & Pressure: - Effect of Temperature (at constant pressure): As T increases, gas particles gain kinetic energy, move faster, and collide with container walls more frequently and with greater force \rightarrow gas expands (volume increases). - Effect of Pressure (at constant temperature): As external P increases, gas particles are forced closer together \rightarrow empty space between them decreases \rightarrow volume decreases. 1.2 Diffusion - Definition: The net movement of particles from a region of higher concentration to a region of lower concentration down a concentration gradient as a result of their random motion. - Kinetic Particle Explanation: Gas/liquid particles move randomly and collide elastically with other particles, causing them to spread out evenly over time. - Effect of Relative Molecular Mass (M_r): - Rate of diffusion is inversely proportional to M_r (\text{Rate} \propto \frac{1}{\sqrt{M_r}}). - Gases with lower M_r diffuse faster than gases with higher M_r at the same temperature. - Apparatus Proof: Cotton wool soaked in concentrated aqueous ammonia (\text{NH}_3, M_r = 17) placed at one end of a glass tube and concentrated hydrochloric acid (\text{HCl}, M_r = 36.5) at the other. A dense white ring of solid ammonium chloride (\text{NH}_4\text{Cl}) forms closer to the \text{HCl} end because \text{NH}_3 is lighter and diffuses faster. SECTION 2: ATOMS, ELEMENTS AND COMPOUNDS 2.1 Elements, compounds and mixtures - Element: A substance made of only one type of atom; cannot be chemically broken down into simpler substances. - Compound: A substance made of two or more different elements chemically combined in fixed proportions. - Mixture: Two or more substances (elements or compounds) physically mixed together but not chemically bonded; can be separated by physical separation techniques. 2.2 Atomic structure and the Periodic Table - Structure of the Atom: Central dense nucleus containing protons and neutrons, surrounded by electrons orbiting in concentric energy levels (shells). - Relative Masses & Charges: | Particle | Relative Charge | Relative Mass | Location | | :----------- | :-------------: | :---------------------------: | :-------------- | | **Proton** | $+1$ | $1$ | Nucleus | | **Neutron** | $0$ | $1$ | Nucleus | | **Electron** | $-1$ | $\frac{1}{1840}$ (negligible) | Electron shells | - Proton number / Atomic number (Z): Number of protons in the nucleus of an atom. - Mass number / Nucleon number (A): Total number of protons and neutrons in the nucleus (A = Z + N). - Electronic Configuration (Elements 1 to 20): - Maximum shell capacities: Shell 1 holds up to 2, Shell 2 up to 8, Shell 3 up to 8, Shell 4 up to 2. - Examples: \text{H} = 1, \text{C} = 2,4, \text{Na} = 2,8,1, \text{Ca} = 2,8,8,2. - Periodic Table Rules: - Group VIII/0 noble gases have a full outer electron shell (duplet for \text{He}, octet for others), making them unreactive. - Number of outer-shell electrons = Group number (Groups I to VII). - Number of occupied electron shells = Period number. 2.3 Isotopes - Definition: Different atoms of the same element that have the same number of protons but different numbers of neutrons. - Notation: ^{A}_{Z}\text{X} (e.g., ^{12}_{6}\text{C} has 6p, 6n, 6e; ^{35}_{17}\text{Cl}^- has 17p, 18n, 18e). - Chemical Properties: Isotopes of the same element have identical chemical properties because they have the same number of electrons and the same electronic configuration. - Relative Atomic Mass (A_r) Calculation: \mathbf{A_r = \frac{\sum (\text{isotopic mass} \times \text{percentage abundance})}{100}} 2.4 Ions and ionic bonds - Ion Formation: - Metals lose outer electrons to form positive ions (cations). - Non-metals gain electrons to form negative ions (anions). - Ionic Bond: A strong electrostatic attraction between oppositely charged ions. - Giant Ionic Lattice Structure: Regular 3D array of alternating positive and negative ions held by omnidirectional electrostatic forces. - Properties Explained by Structure: - High melting and boiling points: Strong electrostatic attraction between oppositely charged ions requires large amounts of thermal energy to break. - Electrical conductivity: Solid ionic compounds do not conduct electricity (ions are fixed in the lattice); molten or aqueous ionic compounds conduct electricity because the lattice breaks down and ions become mobile and free to move. 2.5 Simple molecules and covalent bonds - Covalent Bond: Formed when a pair of electrons is shared between two non-metal atoms, leading to noble gas electronic configurations. - Molecules to Know (Dot-and-Cross): - Single bonds: \text{H}_2, \text{Cl}_2, \text{HCl}, \text{H}_2\text{O}, \text{CH}_4, \text{NH}_3, \text{CH}_3\text{OH}. - Double bonds: \text{O}_2, \text{CO}_2\ (\text{O}=\text{C}=\text{O}), \text{C}_2\text{H}_4. - Triple bonds: \text{N}_2\ (\text{N}\equiv\text{N}). - Properties of Simple Molecular Compounds: - Low melting and boiling points: Weak intermolecular forces between molecules require very little energy to overcome (covalent bonds within molecules are strong and do not break during melting/boiling). - Poor electrical conductivity: No mobile ions or delocalised electrons to carry charge. 2.6 Giant covalent structures 1. Diamond: - Each carbon atom is covalently bonded to 4 other carbon atoms in a rigid tetrahedral network. - Hardest natural substance; very high melting point; electrical insulator (no free electrons). Used in cutting and drilling tools. 2. Graphite: - Each carbon atom is covalently bonded to 3 other carbon atoms in flat hexagonal layers. - One electron per carbon atom is delocalised \rightarrow conducts electricity (used as electrodes). - Layers are held by weak intermolecular forces \rightarrow layers slide over each other easily \rightarrow soft, slippery (used as a lubricant). 3. Silicon(IV) Oxide (\text{SiO}_2 / Quartz): - Giant covalent lattice: each \text{Si} is covalently bonded to 4 oxygen atoms; each \text{O} is covalently bonded to 2 silicon atoms. - Hard, high melting point, electrical insulator (mirrors diamond properties). 2.7 Metallic bonding - Definition: The electrostatic attraction between positive ions in a giant metallic lattice and a sea of delocalised electrons. - Properties Explained: - Good electrical and thermal conductivity: Delocalised electrons are mobile and move freely through the lattice. - Malleability and ductility: Regular layers of positive metal ions can slide over one another without breaking the non-directional metallic bond. SECTION 3: STOICHIOMETRY 3.1 Formulae & Equations - Molecular Formula: The actual number and type of different atoms in one molecule of a compound. - Empirical Formula: The simplest whole number ratio of the different atoms or ions in a compound. - Writing Balanced Equations & Ionic Equations: - Balance atoms on both sides. - Include state symbols: (\text{s}) = solid, (\text{l}) = liquid, (\text{g}) = gas, (\text{aq}) = aqueous solution. - To write ionic equations: split soluble ionic compounds (aq) into separate ions, cancel out spectator ions. 3.2 Relative Masses - Relative Atomic Mass (A_r): The average mass of the isotopes of an element compared to \frac{1}{12}\text{th} of the mass of an atom of ^{12}\text{C}. - Relative Molecular/Formula Mass (M_r): The sum of the relative atomic masses of all atoms present in a molecular formula or formula unit. 3.3 The Mole & Quantitative Chemistry \text{Moles } (n) = \frac{\text{Mass in grams } (m)}{\text{Molar mass } (M_r)} \text{Gas Moles at r.t.p. } (n) = \frac{\text{Volume in dm}^3}{24} = \frac{\text{Volume in cm}^3}{24\,000} \text{Concentration } (c) = \frac{\text{Moles } (n)}{\text{Volume in dm}^3} = \frac{n \times 1000}{\text{Volume in cm}^3} \text{Concentration (g/dm}^3\text{)} = \text{Concentration (mol/dm}^3\text{)} \times M_r \text{Number of particles} = \text{Moles } (n) \times 6.02 \times 10^{23} \text{ (Avogadro Constant)} \%\text{ Yield} = \frac{\text{Actual mass of product}}{\text{Theoretical mass of product}} \times 100 \%\text{ Purity} = \frac{\text{Mass of pure substance}}{\text{Total mass of impure sample}} \times 100 \%\text{ Composition by mass} = \frac{\text{Number of atoms of element } \times A_r}{M_r \text{ of compound}} \times 100 SECTION 4: ELECTROCHEMISTRY 4.1 Electrolysis - Definitions: - Electrolysis: The decomposition of an ionic compound, when molten or in aqueous solution, by the passage of an electric current. - Anode: Positive electrode (+); attracts anions; oxidation occurs. - Cathode: Negative electrode (–); attracts cations; reduction occurs. - Electrolyte: The molten or aqueous ionic liquid undergoing electrolysis. - Charge Transfer: - Electrons flow through the external metal wire circuit from anode to cathode. - Ions migrate through the electrolyte towards the oppositely charged electrodes. - At the cathode: cations gain electrons (reduction). - At the anode: anions lose electrons (oxidation). Products of Electrolysis 1. Molten Binary Salts: - Molten \text{PbBr}_2: - Cathode (–): \text{Pb}^{2+} + 2e^- \rightarrow \text{Pb(l)} (Silvery liquid bead) - Anode (+): 2\text{Br}^- \rightarrow \text{Br}_2\text{(g)} + 2e^- (Brown gas) 2. Concentrated Aqueous Sodium Chloride (\text{NaCl} / Brine): - Cathode (–): 2\text{H}^+ + 2e^- \rightarrow \text{H}_2\text{(g)} (Effervescence) - Anode (+): 2\text{Cl}^- \rightarrow \text{Cl}_2\text{(g)} + 2e^- (Pale yellow-green pungent gas) - Residue: \text{Na}^+ and \text{OH}^- remain \rightarrow aqueous sodium hydroxide (\text{NaOH}). 3. Dilute Sulfuric Acid (\text{H}_2\text{SO}_4): - Cathode (–): 2\text{H}^+ + 2e^- \rightarrow \text{H}_2\text{(g)} (Volume ratio = 2) - Anode (+): 4\text{OH}^- \rightarrow 2\text{H}_2\text{O(l)} + \text{O}_2\text{(g)} + 4e^- (Volume ratio = 1) 4. Aqueous Copper(II) Sulfate (\text{CuSO}_4): - Using Inert Electrodes (Carbon/Platinum): - Cathode (–): \text{Cu}^{2+} + 2e^- \rightarrow \text{Cu(s)} (Pink-brown solid forms) - Anode (+): 4\text{OH}^- \rightarrow 2\text{H}_2\text{O} + \text{O}_2 + 4e^- (Effervescence) - Solution: Blue colour fades as \text{Cu}^{2+} ions are removed; solution becomes acidic (\text{H}_2\text{SO}_4). - Using Copper Electrodes (Purification / Refining): - Anode (+) Impure copper: \text{Cu(s)} \rightarrow \text{Cu}^{2+}\text{(aq)} + 2e^- (Anode dissolves, loses mass) - Cathode (–) Pure copper: \text{Cu}^{2+}\text{(aq)} + 2e^- \rightarrow \text{Cu(s)} (Cathode gains mass) - Solution: Blue colour and \text{Cu}^{2+} concentration remain unchanged. 5. Electroplating: - Object to be coated is placed at the Cathode (–). - Coating metal is placed at the Anode (+). - Electrolyte is a soluble salt containing ions of the coating metal. - Purpose: Prevents corrosion and improves appearance. 4.2 Hydrogen–oxygen fuel cells - Principle: Uses hydrogen and oxygen to produce electricity; water is the only chemical product. 2\text{H}_2\text{(g)} + \text{O}_2\text{(g)} \rightarrow 2\text{H}_2\text{O(l)} - Comparison with Petrol/Gasoline Engines: - Advantages: Produces zero greenhouse gases or toxic pollutants at point of use; highly efficient; fewer moving parts (low friction/wear). - Disadvantages: Hydrogen is difficult and expensive to store safely (requires high pressure or cryogenic tanks); hydrogen is highly flammable; lack of hydrogen filling station infrastructure; hydrogen production often uses fossil fuels. SECTION 5: CHEMICAL ENERGETICS 5.1 Exothermic and endothermic reactions - Exothermic: A reaction that transfers thermal energy to the surroundings \rightarrow temperature of surroundings increases (\Delta H is negative). - Endothermic: A reaction that absorbs thermal energy from the surroundings \rightarrow temperature of surroundings decreases (\Delta H is positive). - Enthalpy Change (\Delta H): The transfer of thermal energy during a chemical reaction. - Activation Energy (E_a): The minimum energy that colliding particles must possess to react. - Reaction Pathway Diagrams: - Vertical arrow from reactants to the highest point of the curve \= \mathbf{E_a}. - Vertical arrow from reactants to products = \mathbf{\Delta H} (pointing down for exothermic; pointing up for endothermic). - Bond Energy Dynamics: - Bond breaking is endothermic (energy must be absorbed). - Bond making is exothermic (energy is released). - In an exothermic reaction, more energy is released when making bonds in the products than is absorbed to break bonds in the reactants. \mathbf{\Delta H = \sum (\text{Bond energies of broken bonds}) - \sum (\text{Bond energies of formed bonds})} SECTION 6: CHEMICAL REACTIONS 6.1 Physical and chemical changes - Physical Change: No new chemical substances formed; easily reversible (e.g., changes of state, dissolving). - Chemical Change: New chemical substances formed with different properties; usually irreversible or difficult to reverse. 6.2 Rate of reaction & Collision Theory - Collision Theory Factors: 1. Concentration (solutions) / Pressure (gases): Increasing concentration/pressure increases the number of particles per unit volume \rightarrow particles collide more frequently (higher collision frequency). 2. Surface Area (solids): Powdering a solid increases the exposed surface area \rightarrow more surface particles available for collision \rightarrow higher collision frequency. 3. Temperature: Increasing temperature increases the kinetic energy of particles \rightarrow particles move faster and collide more frequently AND a significantly higher proportion of colliding particles possess energy \ge E_a (leading to more successful collisions per unit time). 4. Catalyst: Increases rate without being chemically changed or consumed \rightarrow provides an alternative pathway with a lower activation energy (E_a). 6.3 Reversible reactions and equilibrium - Dynamic Equilibrium Criteria: 1. Rate of forward reaction = Rate of reverse reaction. 2. Concentrations of reactants and products are constant (no longer changing). 3. Established only in a closed system. - Le Chatelier’s Principle: - Temperature: - Increase T \rightarrow Equilibrium shifts in the endothermic direction. - Decrease T \rightarrow Equilibrium shifts in the exothermic direction. - Pressure (Gases): - Increase P \rightarrow Shifts to side with fewer gas moles. - Decrease P \rightarrow Shifts to side with more gas moles. - Catalyst: No effect on equilibrium position or yield; only increases the rate at which equilibrium is reached. - Reversible Hydration Tests for Water: - Hydrated copper(II) sulfate \xrightleftharpoons{\Delta} Anhydrous copper(II) sulfate + Water \text{CuSO}_4\cdot5\text{H}_2\text{O(s)} \xrightleftharpoons{} \text{CuSO}_4\text{(s)} + 5\text{H}_2\text{O(l)} \quad (\text{Blue } \xrightleftharpoons{} \text{ White}) - Hydrated cobalt(II) chloride \xrightleftharpoons{\Delta} Anhydrous cobalt(II) chloride + Water \text{CoCl}_2\cdot6\text{H}_2\text{O(s)} \xrightleftharpoons{} \text{CoCl}_2\text{(s)} + 6\text{H}_2\text{O(l)} \quad (\text{Pink } \xrightleftharpoons{} \text{ Blue}) - The Haber Process (Ammonia): \text{N}_2\text{(g)} + 3\text{H}_2\text{(g)} \rightleftharpoons 2\text{NH}_3\text{(g)} \quad \Delta H = -92\text{ kJ/mol} - Sources: \text{N}_2 from fractional distillation of liquid air; \text{H}_2 from methane (natural gas) reacting with steam. - Conditions: 450\ ^\circ\text{C}, 20\,000\text{ kPa} (200\text{ atm}), Iron catalyst. - Explanation of conditions: - Temperature (450\ ^\circ\text{C}): Compromise between higher equilibrium yield (favoured by low T) and acceptable reaction rate (favoured by high T). - Pressure (200\text{ atm}): High pressure favours product side (fewer gas moles, 4 \rightarrow 2) and increases rate, but is limited by the high economic cost and safety hazards of high-pressure equipment. - The Contact Process (Sulfuric Acid Stage 2): 2\text{SO}_2\text{(g)} + \text{O}_2\text{(g)} \rightleftharpoons 2\text{SO}_3\text{(g)} \quad \Delta H = -197\text{ kJ/mol} - Sources: \text{SO}_2 from burning sulfur in air or roasting sulfide ores (e.g. zinc blende \text{ZnS}); \text{O}_2 from air. - Conditions: 450\ ^\circ\text{C}, 200\text{ kPa} (2\text{ atm}), Vanadium(V) oxide (\text{V}_2\text{O}_5) catalyst. - Low pressure (2\text{ atm}) is sufficient because yield is already >98\% at this pressure. 6.4 Redox - Oxidation: Loss of electrons; increase in oxidation number; gain of oxygen. - Reduction: Gain of electrons; decrease in oxidation number; loss of oxygen. - Oxidation Number Rules: - Free uncombined element = 0. - Monatomic ion = charge on ion. - Neutral compound: Sum of oxidation numbers = 0. - Polyatomic ion: Sum of oxidation numbers = charge on ion. - Combined \text{O} = -2 (except peroxides = -1). - Combined \text{H} = +1 (except metal hydrides = -1). - Group I = +1, Group II = +2, \text{F} = -1. - Reagent Confirmation: - Oxidising agent: Oxidises other substances and is itself reduced. Test: turns aqueous potassium iodide (\text{KI}) from colourless to brown (2\text{I}^- \rightarrow \text{I}_2). - Reducing agent: Reduces other substances and is itself oxidised. Test: turns acidified aqueous potassium manganate(VII) from purple to colourless (\text{MnO}_4^- \rightarrow \text{Mn}^{2+}). SECTION 7: ACIDS, BASES AND SALTS 7.1 Characteristic properties of acids and bases - Brønsted–Lowry Definitions: - Acid: Proton (\text{H}^+) donor. - Base: Proton (\text{H}^+) acceptor. - In aqueous solution: acids produce \text{H}^+\text{(aq)} ions; alkalis produce \text{OH}^-\text{(aq)} ions. - Neutralisation Ionic Equation: \mathbf{H^+\text{(aq)} + OH^-\text{(aq)} \rightarrow H_2O\text{(l)}} - Reactions of Acids: - \text{Acid} + \text{Metal} \rightarrow \text{Salt} + \text{Hydrogen} - \text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water} - \text{Acid} + \text{Carbonate} \rightarrow \text{Salt} + \text{Water} + \text{Carbon Dioxide} - Reactions of Bases: - \text{Base} + \text{Ammonium salt} \xrightarrow{\Delta} \text{Salt} + \text{Water} + \text{Ammonia gas} - Acid Strength: - Strong acid: Completely dissociates in aqueous solution (\text{HCl(aq)} \rightarrow \text{H}^+\text{(aq)} + \text{Cl}^-\text{(aq)}). - Weak acid: Partially dissociates in aqueous solution (\text{CH}_3\text{COOH(aq)} \rightleftharpoons \text{H}^+\text{(aq)} + \text{CH}_3\text{COO}^-\text{(aq)}). - Indicators Table: | Indicator | Colour in Acid | Colour in Alkali | Colour at End-point | | :------------------ | :------------: | :--------------: | :-----------------: | | **Litmus** | Red | Blue | Purple | | **Thymolphthalein** | Colourless | Blue | Very pale blue | | **Methyl orange** | Red | Yellow | Orange | 7.2 Oxides - Basic Oxides: Oxides of metals (e.g., \text{CuO}, \text{CaO}) that react with acids to produce a salt and water. - Acidic Oxides: Oxides of non-metals (e.g., \text{SO}_2, \text{CO}_2) that react with bases to produce a salt and water. - Amphoteric Oxides: Oxides that react with both acids and bases to produce a salt and water: - \text{Al}_2\text{O}_3 (Aluminium oxide) and \text{ZnO} (Zinc oxide). - Neutral Oxides: Do not react with acids or bases (\text{CO}, \text{NO}, \text{N}_2\text{O}). 7.3 Preparation of salts - Solubility Rules: - Soluble: All nitrates; all sodium, potassium, and ammonium salts; all chlorides (except \text{Ag}^+, \text{Pb}^{2+}); all sulfates (except \text{Ba}^{2+}, \text{Ca}^{2+}, \text{Pb}^{2+}). - Insoluble: All carbonates and hydroxides (except \text{Na}^+, \text{K}^+, \text{NH}_4^+). \text{Ca(OH)}_2 is partially soluble. - Salt Preparation Methods: 1. Excess Insoluble Reactant Method (Soluble salt from acid + insoluble base/metal/carbonate): - Add excess solid to warm acid until effervescence stops / solid remains unreacted. - Filter to remove excess unreacted solid residue. - Heat filtrate to point of crystallisation (dip cold glass rod to test). - Leave to cool and crystallise slowly. - Filter crystals, wash with small volume of cold distilled water, dry between filter papers. 2. Titration Method (Soluble salt from acid + alkali/soluble carbonate): - Pipette exact volume of alkali into conical flask with indicator. - Titrate using burette with acid until indicator changes colour; note volume. - Repeat without indicator using identical volumes. - Evaporate partially, cool to crystallise, filter, and dry. 3. Precipitation Method (Insoluble salt from two soluble salts): - Mix aqueous solutions of two soluble salts. - Filter the mixture to collect the precipitate as residue. - Wash residue thoroughly with distilled water to remove soluble impurities. - Dry the residue in a warm oven. - Water of Crystallisation: Water molecules chemically bonded into the crystal structure of a salt (e.g., \text{CuSO}_4\cdot5\text{H}_2\text{O}). Hydrated contains water of crystallisation; anhydrous contains no water. SECTION 8: THE PERIODIC TABLE 8.1 Arrangement of elements - Elements are arranged in order of increasing atomic (proton) number. - Across a period: properties change from metallic to non-metallic. - Elements in the same group have the same number of outer-shell electrons and similar chemical properties. 8.2 Group I properties (Alkali Metals: \text{Li, Na, K, Rb, Cs}) - Relatively soft metals (can be cut with a knife); low densities (float on water); low melting points. - Trends Down Group I: - Melting point decreases. - Density increases. - Reactivity increases (outer electron is further from nucleus, more shielded, weaker attraction, lost more easily). - Reaction with water: 2M\text{(s)} + 2\text{H}_2\text{O(l)} \rightarrow 2M\text{OH(aq)} + \text{H}_2\text{(g)} - Produces alkaline solution (\text{pH} > 12). 8.3 Group VII properties (Halogens: \text{Cl, Br, I}) - Diatomic non-metals (\text{Cl}_2, \text{Br}_2, \text{I}_2). - Chlorine: Pale yellow-green gas. - Bromine: Red-brown liquid (forms orange vapor). - Iodine: Grey-black solid (sublimes to purple vapor). - Trends Down Group VII: - Density increases; melting and boiling points increase. - Reactivity decreases (atomic radius increases, greater electron shielding, harder to attract the incoming electron). - Displacement Reactions: - A more reactive halogen displaces a less reactive halide ion from solution: \text{Cl}_2\text{(aq)} + 2\text{Br}^-\text{(aq)} \rightarrow 2\text{Cl}^-\text{(aq)} + \text{Br}_2\text{(aq)} \quad (\text{colourless to orange}) \text{Br}_2\text{(aq)} + 2\text{I}^-\text{(aq)} \rightarrow 2\text{Br}^-\text{(aq)} + \text{I}_2\text{(aq)} \quad (\text{orange to brown}) 8.4 Transition elements - Typical metallic properties: High densities, high melting points, hard and strong. - Characteristic chemical properties: 1. Form coloured compounds. 2. Have ions with variable oxidation numbers (e.g., \text{Fe}^{2+} and \text{Fe}^{3+}). 3. Act as catalysts as elements and in compounds. 8.5 Noble gases (Group VIII / 0) - Unreactive, monatomic gases. - Unreactive because they have a full outer shell of electrons (stable octet/duplet). SECTION 9: METALS 9.1 Properties of metals - Physical: Good electrical and thermal conductors; malleable and ductile; high melting and boiling points (except Group I and mercury). - Chemical: React with oxygen to form basic oxides; react with acids to form salt and hydrogen. 9.2 Uses of metals - Aluminium: - Aircraft bodies: Low density, high strength-to-weight ratio. - Overhead power cables: Low density and good electrical conductivity. - Food containers: Resistance to corrosion (protective oxide layer). - Copper: Electrical wiring: Excellent electrical conductivity and ductility. 9.3 Alloys and their properties - Alloy: A mixture of a metal with other elements (e.g. brass \= \text{copper} + \text{zinc}; stainless steel \= \text{iron} + \text{chromium} + \text{nickel} + \text{carbon}). - Why Alloys are Harder than Pure Metals: Pure metals consist of regular layers of identical atoms that slide over each other easily. Alloys contain different sized atoms which distort the regular lattice layers, preventing them from sliding over one another easily. 9.4 Reactivity series \mathbf{K > Na > Ca > Mg > Al > [C] > Zn > Fe > [H] > Cu > Ag > Au} - Reactions: - With cold water: \text{K, Na, Ca} react vigorously to form metal hydroxide and \text{H}_2. - With steam: \text{Mg} reacts vigorously to form metal oxide and \text{H}_2: \text{Mg(s)} + \text{H}_2\text{O(g)} \rightarrow \text{MgO(s)} + \text{H}_2\text{(g)} - With dilute \text{HCl}: Metals above \text{H} react to form salt and \text{H}_2; rate depends on reactivity. \text{Cu, Ag, Au} do not react. - Apparent Unreactivity of Aluminium: Aluminium is high in the reactivity series but does not react with water/acid under normal conditions due to a tough, adhering, impermeable protective oxide layer (\text{Al}_2\text{O}_3). 9.5 Corrosion of metals - Rusting: Requires water AND oxygen to form hydrated iron(III) oxide (\text{Fe}_2\text{O}_3\cdot x\text{H}_2\text{O}). - Barrier Methods: Painting, greasing, plastic coating (prevent contact with air and water). - Galvanising & Sacrificial Protection: Coating iron with zinc. Zinc acts as a physical barrier and provides sacrificial protection: zinc is more reactive than iron, so zinc oxidises/loses electrons preferentially (\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-), preventing iron from oxidising even if scratched. 9.6 Extraction of metals - Ease of extraction depends on reactivity: unreactive metals occur native (\text{Au, Ag}); moderate reactivity metals reduced by carbon; highly reactive metals extracted by electrolysis. - Blast Furnace (Iron Extraction): - Ore: Hematite (\text{Fe}_2\text{O}_3). - Equations: 1. \text{C} + \text{O}_2 \rightarrow \text{CO}_2 (coke burns to provide heat) 2. \text{C} + \text{CO}_2 \rightarrow 2\text{CO} (carbon dioxide reduced to carbon monoxide) 3. \text{Fe}_2\text{O}_3 + 3\text{CO} \rightarrow 2\text{Fe} + 3\text{CO}_2 (iron(III) oxide reduced by carbon monoxide) 4. \text{CaCO}_3 \rightarrow \text{CaO} + \text{CO}_2 (thermal decomposition of limestone) 5. \text{CaO} + \text{SiO}_2 \rightarrow \text{CaSiO}_3 (calcium oxide neutralises acidic silicon(IV) oxide impurity to form slag) - Extraction of Aluminium: - Ore: Bauxite (purified to \text{Al}_2\text{O}_3). - Role of Cryolite: Dissolves aluminium oxide, lowering melting point from 2000\ ^\circ\text{C} to \sim 950\ ^\circ\text{C} and improving electrical conductivity. - Cathode (–) Carbon: \text{Al}^{3+} + 3e^- \rightarrow \text{Al(l)} - Anode (+) Carbon: 2\text{O}^{2-} \rightarrow \text{O}_2\text{(g)} + 4e^- - Carbon anodes react with oxygen gas at high temperatures (\text{C} + \text{O}_2 \rightarrow \text{CO}_2), burning away and requiring regular replacement. SECTION 10: CHEMISTRY OF THE ENVIRONMENT 10.1 Water - Chemical Tests for Water: - Anhydrous cobalt(II) chloride paper turns from blue to pink. - Anhydrous copper(II) sulfate turns from white to blue. - Purity Test: Pure water freezes at exactly 0\ ^\circ\text{C} and boils at exactly 100\ ^\circ\text{C} at 1\text{ atm}. - Domestic Water Treatment: 1. Sedimentation and filtration through sand/gravel beds to remove insoluble solid matter. 2. Activated carbon beds to remove unpleasant tastes and odours. 3. Chlorination to kill harmful bacteria and microbes. 10.2 Fertilisers - NPK Fertilisers: Provide three essential elements for plant growth: Nitrogen (N) (for chlorophyll/proteins), Phosphorus (P) (for root development), and Potassium (K) (for flower/fruit growth). - Ammonium salts and nitrates are common sources of nitrogen. 10.3 Air quality and climate - Clean, Dry Air: 78\% \text{ N}_2, 21\% \text{ O}_2, remainder noble gases and 0.04\% \text{ CO}_2. - Pollutants: - \text{CO}_2: Complete combustion of fossil fuels \rightarrow greenhouse gas causing global warming. - \text{CO}: Incomplete combustion of fossil fuels \rightarrow toxic gas (binds to haemoglobin). - Particulates (soot): Incomplete combustion \rightarrow respiratory problems and cancer. - \text{CH}_4: Decomposition of vegetation, animal digestion \rightarrow greenhouse gas. - \text{NO}_x (\text{NO, NO}_2): Formed in car engines at high temperatures \rightarrow acid rain, photochemical smog, respiratory problems. - \text{SO}_2: Combustion of fossil fuels containing sulfur \rightarrow acid rain. - Greenhouse Effect: Greenhouse gases (\text{CO}_2, \text{CH}_4) absorb thermal energy radiated from the Earth's surface and re-emit it in all directions, reducing thermal energy loss to space. - Catalytic Converters in Vehicles: 2\text{CO} + 2\text{NO} \rightarrow 2\text{CO}_2 + \text{N}_2 - Photosynthesis: 6\text{CO}_2 + 6\text{H}_2\text{O} \xrightarrow{\text{light, chlorophyll}} \text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2 SECTION 11: ORGANIC CHEMISTRY 11.1 Formulae, functional groups and terminology - Hydrocarbon: A compound containing hydrogen and carbon ONLY. - General Formulae: - Alkanes: \text{C}_n\text{H}_{2n+2} - Alkenes: \text{C}_n\text{H}_{2n} - Alcohols: \text{C}_n\text{H}_{2n+1}\text{OH} - Carboxylic acids: \text{C}_n\text{H}_{2n+1}\text{COOH} - Structural Isomers: Compounds with the same molecular formula but different structural formulae. - Example: \text{C}_4\text{H}_{10} exists as butane (\text{CH}_3\text{CH}_2\text{CH}_2\text{CH}_3) and 2-methylpropane (\text{CH}_3\text{CH(CH}_3)\text{CH}_3). - Homologous Series Characteristics: Same functional group; same general formula; successive members differ by a -\text{CH}_2- unit; gradual trend in physical properties; similar chemical properties. 11.2 Naming organic compounds (IUPAC) - Root: 1\text{C} = meth-, 2\text{C} = eth-, 3\text{C} = prop-, 4\text{C} = but-. - Suffix: alkane (-ane), alkene (-ene), alcohol (-ol), carboxylic acid (-oic acid), ester (-yl ... -oate). - Specific isomers to draw/name: - But-1-ene (\text{CH}_2=\text{CHCH}_2\text{CH}_3) vs But-2-ene (\text{CH}_3\text{CH}=\text{CHCH}_3) - Propan-1-ol vs Propan-2-ol; Butan-1-ol vs Butan-2-ol. 11.3 Fuels & Fractional Distillation - Petroleum is a mixture of hydrocarbons separated by fractional distillation based on boiling points. - Fractions from top to bottom: 1. Refinery gas: Bottled gas for heating and cooking. 2. Gasoline (petrol): Fuel for cars. 3. Naphtha: Chemical feedstock for plastics. 4. Kerosene (paraffin): Jet fuel. 5. Diesel oil (gas oil): Fuel for diesel engines. 6. Fuel oil: Fuel for ships and home heating. 7. Lubricating oil: Lubricants, waxes, polishes. 8. Bitumen: Road surfacing. - Trends from bottom to top: Chain length decreases; volatility increases; boiling points decrease; viscosity decreases (flows more easily). 11.4 Alkanes - Saturated hydrocarbons (only single \text{C}-\text{C} bonds). Generally unreactive except combustion and substitution. - Photochemical Substitution: \text{CH}_4 + \text{Cl}_2 \xrightarrow{\text{UV light}} \text{CH}_3\text{Cl} + \text{HCl} - UV light provides the activation energy (E_a). 11.5 Alkenes - Unsaturated hydrocarbons with a \text{C}=\text{C} double bond. - Manufacture by Cracking: Large alkanes heated with catalyst (silica/alumina/zeolite) to form smaller alkanes and alkenes: \text{C}_{10}\text{H}_{22} \rightarrow \text{C}_8\text{H}_{18} + \text{C}_2\text{H}_4 - Addition Reactions: Form only one product. 1. Aqueous Bromine Test: Turns from orange/brown to colourless (decolourises). \text{C}_2\text{H}_4 + \text{Br}_2 \rightarrow \text{CH}_2\text{BrCH}_2\text{Br} \quad (\text{1,2-dibromoethane}) 2. Hydrogenation: \text{C}_2\text{H}_4 + \text{H}_2 \xrightarrow{\text{Ni catalyst, } 150-200\ ^\circ\text{C}} \text{C}_2\text{H}_6 3. Hydration: \text{C}_2\text{H}_4\text{(g)} + \text{H}_2\text{O(g)} \xrightarrow{\text{conc. } \text{H}_3\text{PO}_4, 300\ ^\circ\text{C}, 60\text{ atm}} \text{C}_2\text{H}_5\text{OH(g)} 11.6 Alcohols - Manufacture of Ethanol: - Fermentation: Aqueous glucose, yeast, 25-35\ ^\circ\text{C}, anaerobic (absence of \text{O}_2). Renewable batch process; slow; yields dilute ethanol. - Catalytic Hydration of Ethene: Continuous; fast; yields pure ethanol; relies on non-renewable petroleum. - Combustion: \text{C}_2\text{H}_5\text{OH} + 3\text{O}_2 \rightarrow 2\text{CO}_2 + 3\text{H}_2\text{O} 11.7 Carboxylic acids - Weak acids containing the -\text{COOH} functional group. - Formation by Oxidation of Ethanol: - With acidified aqueous potassium manganate(VII) under heat. - Bacterial oxidation during vinegar production. - Esterification: \text{Carboxylic acid} + \text{Alcohol} \xrightleftharpoons{\text{acid catalyst (conc. } \text{H}_2\text{SO}_4\text{)}} \text{Ester} + \text{Water} - Example: \text{CH}_3\text{COOH} + \text{CH}_3\text{OH} \rightleftharpoons \text{CH}_3\text{COOCH}_3 + \text{H}_2\text{O} (methyl ethanoate). 11.8 Polymers 1. Addition Polymerisation: Double bond breaks; unsaturated monomers link to form long-chain saturated polymers (e.g. ethene \rightarrow poly(ethene)). Only one product forms. 2. Condensation Polymerisation: Monomers link with the elimination of a small molecule (usually \text{H}_2\text{O}). - Polyamides (Nylon): Formed from dicarboxylic acid and diamine monomers; linked by amide linkages (-\text{CONH}-). - Polyesters (PET): Formed from dicarboxylic acid and diol monomers; linked by ester linkages (-\text{COO}-). PET can be chemically hydrolysed back into monomers and re-polymerised. - Proteins: Natural polyamides formed from amino acid monomers with general structure \text{H}_2\text{N}-\text{CH(R)}-\text{COOH}; linked by amide/peptide linkages. 3. Plastics & Environment: Non-biodegradable; accumulation in landfill and oceans; toxic gas formation during incineration (\text{CO}, \text{HCl}, \text{HCN}). SECTION 12: EXPERIMENTAL TECHNIQUES & CHEMICAL ANALYSIS 12.1 Experimental design & apparatus - Precision Tools: - Time: Stop-watch (accuracy of 1\text{ s}). - Temperature: Thermometer (graduations to 1\ ^\circ\text{C}). - Mass: Digital balance (precision of at least 0.1\text{ g}). - Volume: Volumetric pipette (fixed volume, e.g. 25.0\text{ cm}^3), burette (variable volume to \pm 0.05\text{ cm}^3), measuring cylinder (approximate volume), gas syringe (gas volumes up to 100\text{ cm}^3). - Definitions: - Solute: A substance that is dissolved in a solvent. - Solvent: A substance that dissolves a solute. - Solution: A mixture of one or more solutes dissolved in a solvent. - Saturated solution: A solution containing the maximum concentration of a solute dissolved in the solvent at a specified temperature. - Residue: Substance remaining after filtration, evaporation, or distillation. - Filtrate: Liquid or solution that passes through a filter paper. 12.2 Acid–base titrations - Rinse burette with the solution it will hold. - Use a volumetric pipette with filler to measure exactly 25.0\text{ cm}^3 of solution into a conical flask. - Place conical flask on a white tile (to clearly see the indicator colour change). - Swirl flask continuously while adding titrant from burette dropwise near the end-point. 12.3 Chromatography - Draw baseline in pencil (pencil graphite does not dissolve in solvent). - Ensure solvent level in beaker is below the pencil baseline (prevents spots dissolving off paper into solvent). - Use a locating agent (e.g. ninhydrin) to visualise colourless spots (e.g. amino acids). R_f = \frac{\text{Distance travelled by substance}}{\text{Distance travelled by solvent front}} 12.5 Qualitative Analysis Master Tables A. Tests for Anions | Anion | Test | Test Result | | :--------------------------------- | :------------------------------------------------------------------ | :---------------------------------------------------------------- | | **Carbonate ($\text{CO}_3^{2-}$)** | Add dilute acid, test gas with limewater | Effervescence; gas turns **limewater milky** ($\text{CO}_2$) | | **Chloride ($\text{Cl}^-$)** | Acidify with dilute $\text{HNO}_3$, add aqueous $\text{AgNO}_3$ | **White precipitate** | | **Bromide ($\text{Br}^-$)** | Acidify with dilute $\text{HNO}_3$, add aqueous $\text{AgNO}_3$ | **Cream precipitate** | | **Iodide ($\text{I}^-$)** | Acidify with dilute $\text{HNO}_3$, add aqueous $\text{AgNO}_3$ | **Yellow precipitate** | | **Nitrate ($\text{NO}_3^-$)** | Add aqueous $\text{NaOH}$, aluminium foil; warm gently | Gas produced turns **damp red litmus paper blue** ($\text{NH}_3$) | | **Sulfate ($\text{SO}_4^{2-}$)** | Acidify with dilute $\text{HNO}_3$, add aqueous $\text{Ba(NO}_3)_2$ | **White precipitate** | | **Sulfite ($\text{SO}_3^{2-}$)** | Add acidified aqueous potassium manganate(VII) | Turns from **purple to colourless** | B. Tests for Aqueous Cations | Cation | Effect of Aqueous $\text{NaOH}$ | Effect of Aqueous $\text{NH}_3$ | | :----------------------------------- | :---------------------------------------------------------- | :---------------------------------------------------------- | | **Aluminium ($\text{Al}^{3+}$)** | **White ppt**, soluble in excess (colourless solution) | **White ppt**, insoluble in excess | | **Ammonium ($\text{NH}_4^+$)** | Ammonia produced on warming (turns damp red litmus blue) | — | | **Calcium ($\text{Ca}^{2+}$)** | **White ppt**, insoluble in excess | No ppt (or very slight white ppt) | | **Chromium(III) ($\text{Cr}^{3+}$)** | **Green ppt**, soluble in excess (green solution) | **Green ppt**, insoluble in excess | | **Copper(II) ($\text{Cu}^{2+}$)** | **Light blue ppt**, insoluble in excess | **Light blue ppt**, soluble in excess (dark blue solution) | | **Iron(II) ($\text{Fe}^{2+}$)** | **Green ppt**, insoluble in excess; turns brown on standing | **Green ppt**, insoluble in excess; turns brown on standing | | **Iron(III) ($\text{Fe}^{3+}$)** | **Red-brown ppt**, insoluble in excess | **Red-brown ppt**, insoluble in excess | | **Zinc ($\text{Zn}^{2+}$)** | **White ppt**, soluble in excess (colourless solution) | **White ppt**, soluble in excess (colourless solution) | C. Tests for Gases | Gas | Test and Test Result | | :--------------------------------- | :----------------------------------------------------------------------------- | | **Ammonia ($\text{NH}_3$)** | Turns **damp red litmus paper blue** | | **Carbon dioxide ($\text{CO}_2$)** | Turns **limewater milky** | | **Chlorine ($\text{Cl}_2$)** | **Bleaches damp litmus paper** | | **Hydrogen ($\text{H}_2$)** | **'Pops' with a lighted splint** | | **Oxygen ($\text{O}_2$)** | **Relights a glowing splint** | | **Sulfur dioxide ($\text{SO}_2$)** | Turns acidified aqueous potassium manganate(VII) from **purple to colourless** | D. Flame Tests for Metal Ions | Metal Ion | Flame Colour | | :-------------------------------- | :-------------- | | **Lithium ($\text{Li}^+$)** | **Red** | | **Sodium ($\text{Na}^+$)** | **Yellow** | | **Potassium ($\text{K}^+$)** | **Lilac** | | **Calcium ($\text{Ca}^{2+}$)** | **Orange-red** | | **Barium ($\text{Ba}^{2+}$)** | **Light green** | | **Copper(II) ($\text{Cu}^{2+}$)** | **Blue-green** |