IGCSE Chemistry (0620) — Complete Revision Notes (Core & Extended)

<h1 class="notes-h1">IGCSE Chemistry (0620) — Complete Revision Notes</h1>
<h2 class="notes-h2">Cambridge Assessment International Education</h2>
<h3 class="notes-h3">Syllabus Code: 0620 | Core & Extended</h3>
<hr class="section-divider">
<p><strong>Prepared by:</strong> CBC Edu Kenya | cbcedukenya.com</p>
<p><strong>Syllabus Version:</strong> 0620 (2023–2025 and 2026 onwards)</p>
<p><strong>Coverage:</strong> All topics — Core and Extended</p>
<p><strong>Note:</strong> Original revision notes aligned to the Cambridge IGCSE Chemistry (0620) syllabus. Not official Cambridge materials.</p>
<hr class="section-divider">
<h2 class="notes-h2">PAPER OVERVIEW</h2>
<div class="table-wrap"><table class="notes-table">
<thead>
<tr><th>Paper</th><th>Duration</th><th>Marks</th><th>Notes</th></tr>
</thead><tbody>
<tr><td>Paper 1 (Core MCQ)</td><td>45 min</td><td>40</td><td>40 MCQs</td></tr>
<tr><td>Paper 2 (Core Structured)</td><td>1 hr 15 min</td><td>80</td><td>Short + structured</td></tr>
<tr><td>Paper 3 (Extended MCQ)</td><td>45 min</td><td>40</td><td>40 MCQs (harder)</td></tr>
<tr><td>Paper 4 (Extended Structured)</td><td>1 hr 15 min</td><td>80</td><td>Short + structured</td></tr>
<tr><td>Paper 5 (Practical)</td><td>1 hr 15 min</td><td>40</td><td>OR</td></tr>
<tr><td>Paper 6 (Alternative to Practical)</td><td>1 hr</td><td>40</td><td>Written</td></tr>
</tbody></table></div>
<hr class="section-divider">
<h2 class="notes-h2">TOPIC 1: STATES OF MATTER</h2>
<h3 class="notes-h3">1.1 Kinetic Theory</h3>
<div class="table-wrap"><table class="notes-table">
<thead>
<tr><th>State</th><th>Particle arrangement</th><th>Movement</th><th>Forces</th></tr>
</thead><tbody>
<tr><td>Solid</td><td>Regular, close-packed lattice</td><td>Vibrate about fixed positions</td><td>Strong</td></tr>
<tr><td>Liquid</td><td>Close, random arrangement</td><td>Move freely but limited</td><td>Moderate</td></tr>
<tr><td>Gas</td><td>Widely spread, random</td><td>Fast, in all directions</td><td>Negligible</td></tr>
</tbody></table></div>
<p><strong>State changes:</strong></p>
<ul class="notes-list">
<li>Solid → Liquid: <strong>melting</strong> (particles gain energy, overcome lattice forces)</li>
<li>Liquid → Gas: <strong>evaporation/boiling</strong> (particles gain energy, escape liquid)</li>
<li>Gas → Liquid: <strong>condensation</strong></li>
<li>Liquid → Solid: <strong>freezing</strong></li>
<li>Solid → Gas: <strong>sublimation</strong> (iodine, dry ice)</li>
</ul>
<h3 class="notes-h3">1.2 Diffusion</h3>
<ul class="notes-list">
<li>Random movement of particles from high to low concentration</li>
<li>Faster in gases than in liquids (more kinetic energy)</li>
<li>Faster at higher temperatures</li>
</ul>
<hr class="section-divider">
<h2 class="notes-h2">TOPIC 2: ATOMS AND THE PERIODIC TABLE</h2>
<h3 class="notes-h3">2.1 Atomic Structure</h3>
<p><strong>Subatomic particles:</strong></p>
<div class="table-wrap"><table class="notes-table">
<thead>
<tr><th>Particle</th><th>Location</th><th>Relative mass</th><th>Relative charge</th></tr>
</thead><tbody>
<tr><td>Proton</td><td>Nucleus</td><td>1</td><td>+1</td></tr>
<tr><td>Neutron</td><td>Nucleus</td><td>1</td><td>0</td></tr>
<tr><td>Electron</td><td>Shells around nucleus</td><td>1/1840 (negligible)</td><td>-1</td></tr>
</tbody></table></div>
<p><strong>Key definitions:</strong></p>
<ul class="notes-list">
<li><strong>Atomic number (Z):</strong> Number of protons (= number of electrons in a neutral atom)</li>
<li><strong>Mass number (A):</strong> Number of protons + neutrons</li>
<li><strong>Number of neutrons:</strong> A − Z</li>
<li><strong>Isotopes:</strong> Atoms of the same element with the same atomic number but different mass numbers (different numbers of neutrons)</li>
</ul>
<p><strong>Worked Example:</strong> ⁶³₂₉Cu has:</p>
<ul class="notes-list">
<li>29 protons, 29 electrons, 63 − 29 = <strong>34 neutrons</strong></li>
</ul>
<h3 class="notes-h3">2.2 Electronic Configuration</h3>
<p>Electrons occupy shells (energy levels):</p>
<ul class="notes-list">
<li>Shell 1: maximum 2 electrons</li>
<li>Shell 2: maximum 8 electrons</li>
<li>Shell 3: maximum 8 electrons (in this course)</li>
</ul>
<p><strong>Examples:</strong></p>
<div class="table-wrap"><table class="notes-table">
<thead>
<tr><th>Element</th><th>Atomic number</th><th>Electronic configuration</th></tr>
</thead><tbody>
<tr><td>Hydrogen</td><td>1</td><td>1</td></tr>
<tr><td>Carbon</td><td>6</td><td>2,4</td></tr>
<tr><td>Sodium</td><td>11</td><td>2,8,1</td></tr>
<tr><td>Chlorine</td><td>17</td><td>2,8,7</td></tr>
<tr><td>Calcium</td><td>20</td><td>2,8,8,2</td></tr>
</tbody></table></div>
<h3 class="notes-h3">2.3 The Periodic Table</h3>
<p><strong>Groups</strong> (vertical columns): Elements with the same number of outer shell electrons (same chemical properties)</p>
<p><strong>Periods</strong> (horizontal rows): Elements with electrons in the same number of shells</p>
<p><strong>Key groups:</strong></p>
<div class="table-wrap"><table class="notes-table">
<thead>
<tr><th>Group</th><th>Name</th><th>Key property</th></tr>
</thead><tbody>
<tr><td>1</td><td>Alkali metals</td><td>1 outer electron; reactive metals</td></tr>
<tr><td>2</td><td>Alkaline earth metals</td><td>2 outer electrons</td></tr>
<tr><td>7 (17)</td><td>Halogens</td><td>7 outer electrons; reactive non-metals</td></tr>
<tr><td>0 (18)</td><td>Noble gases</td><td>Full outer shell; very unreactive</td></tr>
<tr><td>Transition metals</td><td>(middle block)</td><td>Variable valency; coloured compounds; catalysts</td></tr>
</tbody></table></div>
<p><strong>Trends across a period (left to right):</strong></p>
<ul class="notes-list">
<li>Metallic character decreases</li>
<li>Non-metallic character increases</li>
<li>Atomic radius decreases</li>
</ul>
<h3 class="notes-h3">2.4 Group 1 — Alkali Metals</h3>
<p>Li, Na, K, Rb, Cs, Fr</p>
<p><strong>Properties:</strong></p>
<ul class="notes-list">
<li>Soft, shiny metals (can be cut with a knife)</li>
<li>React with water → metal hydroxide + hydrogen</li>
<li>2Na + 2H₂O → 2NaOH + H₂</li>
<li>Reactivity <strong>increases</strong> down the group (outer electron in higher, more shielded shell → easier to lose)</li>
<li>K burns with lilac/violet flame; Na with yellow; Li with red</li>
</ul>
<p><strong>Trend down the group:</strong> More reactive, lower melting point, lower density</p>
<h3 class="notes-h3">2.5 Group 7 — Halogens</h3>
<p>F, Cl, Br, I, At</p>
<p><strong>Properties:</strong></p>
<ul class="notes-list">
<li>Exist as diatomic molecules (F₂, Cl₂, Br₂, I₂)</li>
<li>React with metals → metal halides (ionic compounds)</li>
<li>2Fe + 3Cl₂ → 2FeCl₃</li>
<li>React with hydrogen → hydrogen halides</li>
<li>H₂ + Cl₂ → 2HCl</li>
<li>Reactivity <strong>decreases</strong> down the group (harder to gain electron as atomic radius increases)</li>
</ul>
<p><strong>Displacement reactions:</strong></p>
<ul class="notes-list">
<li>A more reactive halogen displaces a less reactive one from its salt</li>
<li>Cl₂ + 2KBr → 2KCl + Br₂ (chlorine displaces bromine)</li>
<li>Cl₂ + 2KI → 2KCl + I₂ (chlorine displaces iodine)</li>
<li>Br₂ + KI → KBr + ½I₂ (bromine displaces iodine)</li>
</ul>
<p><strong>Colours:</strong></p>
<ul class="notes-list">
<li>F₂: pale yellow gas</li>
<li>Cl₂: yellow-green gas</li>
<li>Br₂: red-brown liquid/vapour</li>
<li>I₂: grey/purple solid; purple vapour</li>
</ul>
<h3 class="notes-h3">2.6 Noble Gases (Group 0/18)</h3>
<p>He, Ne, Ar, Kr, Xe, Rn</p>
<ul class="notes-list">
<li>Full outer electron shells → very stable → almost no reactions</li>
<li>Uses: inert atmospheres (welding), balloons (He), lighting (Ne, Kr, Xe)</li>
</ul>
<hr class="section-divider">
<h2 class="notes-h2">TOPIC 3: CHEMICAL BONDING</h2>
<h3 class="notes-h3">3.1 Ionic Bonding</h3>
<ul class="notes-list">
<li>Occurs between metals and non-metals</li>
<li>Metal <strong>loses</strong> electrons → cation (positive ion)</li>
<li>Non-metal <strong>gains</strong> electrons → anion (negative ion)</li>
<li><strong>Strong electrostatic attraction</strong> between oppositely charged ions</li>
<li>Forms a <strong>giant ionic lattice</strong></li>
</ul>
<p><strong>Properties of ionic compounds:</strong></p>
<ul class="notes-list">
<li>High melting/boiling points</li>
<li>Hard but brittle (ions repel when shifted)</li>
<li>Conduct electricity when <strong>dissolved in water or molten</strong> (ions free to move), NOT when solid (ions fixed in lattice)</li>
<li>Usually soluble in water</li>
</ul>
<p><strong>Worked Example:</strong> Formation of NaCl:</p>
<ul class="notes-list">
<li>Na (2,8,1) → loses 1 electron → Na⁺ (2,8) + e⁻</li>
<li>Cl (2,8,7) + e⁻ → Cl⁻ (2,8,8)</li>
</ul>
<h3 class="notes-h3">3.2 Covalent Bonding</h3>
<ul class="notes-list">
<li>Occurs between non-metals</li>
<li>Atoms <strong>share</strong> electron pairs</li>
<li>Each shared pair = one covalent bond</li>
<li>Results in discrete molecules or giant covalent structures</li>
</ul>
<p><strong>Common molecules:</strong></p>
<div class="table-wrap"><table class="notes-table">
<thead>
<tr><th>Molecule</th><th>Bonds</th><th>Structure</th></tr>
</thead><tbody>
<tr><td>H₂</td><td>1 single (H-H)</td><td>H:H</td></tr>
<tr><td>Cl₂</td><td>1 single</td><td>Cl:Cl</td></tr>
<tr><td>O₂</td><td>1 double</td><td>O::O</td></tr>
<tr><td>N₂</td><td>1 triple</td><td>N:::N</td></tr>
<tr><td>H₂O</td><td>2 single</td><td>Bent molecule</td></tr>
<tr><td>CO₂</td><td>2 double</td><td>Linear: O=C=O</td></tr>
<tr><td>CH₄</td><td>4 single</td><td>Tetrahedral</td></tr>
<tr><td>NH₃</td><td>3 single</td><td>Pyramidal</td></tr>
</tbody></table></div>
<p><strong>Properties of simple covalent molecules:</strong></p>
<ul class="notes-list">
<li>Low melting/boiling points (weak intermolecular forces between molecules)</li>
<li>Do NOT conduct electricity (no free electrons or ions)</li>
<li>Often insoluble in water (depends on polarity)</li>
</ul>
<p><strong>Giant covalent structures [E]:</strong></p>
<ul class="notes-list">
<li>Diamond: carbon; each C bonded to 4 others in 3D lattice; very hard; very high mp; does not conduct</li>
<li>Graphite: carbon; layers of hexagonal rings; each C bonded to 3 others; layers slide → lubricant; conducts electricity (delocalised electrons)</li>
<li>Silicon dioxide (SiO₂): hard; high mp; does not conduct</li>
</ul>
<h3 class="notes-h3">3.3 Metallic Bonding [E]</h3>
<ul class="notes-list">
<li>Metal cations in a lattice surrounded by a "sea" of delocalised electrons</li>
<li>Electrostatic attraction between cations and electrons = metallic bond</li>
</ul>
<p><strong>Properties explained:</strong></p>
<ul class="notes-list">
<li>High melting points: strong metallic bonds</li>
<li>Good conductors of heat and electricity: free (delocalised) electrons</li>
<li>Malleable/ductile: layers of ions slide without breaking the "sea" of electrons</li>
</ul>
<hr class="section-divider">
<h2 class="notes-h2">TOPIC 4: STOICHIOMETRY</h2>
<h3 class="notes-h3">4.1 Relative Atomic Mass and Relative Molecular Mass</h3>
<ul class="notes-list">
<li><strong>Relative atomic mass (Aᵣ):</strong> Average mass of an atom relative to 1/12 the mass of ¹²C</li>
<li><strong>Relative molecular mass (Mᵣ):</strong> Sum of relative atomic masses in one molecule</li>
<li>Mᵣ of H₂SO₄ = 2(1) + 32 + 4(16) = 98</li>
</ul>
<h3 class="notes-h3">4.2 The Mole</h3>
<ul class="notes-list">
<li>1 mole = 6.02 × 10²³ particles (<strong>Avogadro's number</strong>)</li>
<li>Molar mass = Mᵣ in grams per mole</li>
</ul>
<p><strong>Key formulae:</strong></p>
<ul class="notes-list">
<li>n (moles) = mass (g) / Mᵣ</li>
<li>n = volume (dm³) × concentration (mol/dm³)</li>
<li>n (gas at RTP) = volume (dm³) / 24 (1 mole of gas = 24 dm³ at RTP)</li>
</ul>
<p><strong>Worked Example:</strong> How many moles are in 11 g of CO₂?</p>
<ul class="notes-list">
<li>Mᵣ of CO₂ = 12 + 32 = 44</li>
<li>n = 11/44 = <strong>0.25 mol</strong></li>
</ul>
<h3 class="notes-h3">4.3 Empirical and Molecular Formulae [E]</h3>
<p><strong>Empirical formula:</strong> Simplest whole-number ratio of atoms.</p>
<p><strong>Molecular formula:</strong> Actual number of atoms in one molecule.</p>
<p><strong>Finding empirical formula from % composition:</strong></p>
<ol class="notes-list">
<li>Divide each % by Aᵣ to get molar ratio</li>
<li>Divide all by the smallest value</li>
<li>Round to nearest whole number (multiply if needed)</li>
</ol>
<p><strong>Worked Example:</strong> A compound contains 40% C, 6.7% H, 53.3% O by mass. Find empirical formula.</p>
<ul class="notes-list">
<li>C: 40/12 = 3.33; H: 6.7/1 = 6.7; O: 53.3/16 = 3.33</li>
<li>Ratio C:H:O = 3.33:6.7:3.33 = 1:2:1</li>
<li>Empirical formula: <strong>CH₂O</strong></li>
</ul>
<h3 class="notes-h3">4.4 Chemical Equations</h3>
<p><strong>Balancing equations:</strong> Atoms of each element must be equal on both sides.</p>
<ul class="notes-list">
<li>CH₄ + 2O₂ → CO₂ + 2H₂O ✓</li>
</ul>
<p><strong>Ionic equations [E]:</strong> Show only the particles that change.</p>
<ul class="notes-list">
<li>Full: HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)</li>
<li>Ionic: H⁺(aq) + OH⁻(aq) → H₂O(l)</li>
</ul>
<h3 class="notes-h3">4.5 Reacting Masses Calculations</h3>
<p><strong>Steps:</strong></p>
<ol class="notes-list">
<li>Write and balance the equation</li>
<li>Find moles of the given substance</li>
<li>Use molar ratio to find moles of the unknown</li>
<li>Convert moles to grams (or volume)</li>
</ol>
<p><strong>Worked Example:</strong> What mass of CaCO₃ is needed to produce 4.4 g of CO₂?</p>
<ul class="notes-list">
<li>CaCO₃ → CaO + CO₂</li>
<li>Moles of CO₂ = 4.4/44 = 0.1 mol</li>
<li>Molar ratio 1:1, so moles of CaCO₃ = 0.1 mol</li>
<li>Mass of CaCO₃ = 0.1 × 100 = <strong>10 g</strong></li>
</ul>
<hr class="section-divider">
<h2 class="notes-h2">TOPIC 5: ELECTROCHEMISTRY</h2>
<h3 class="notes-h3">5.1 Electrolysis</h3>
<p><strong>Electrolysis:</strong> Decomposition of a compound using electricity.</p>
<ul class="notes-list">
<li><strong>Electrolyte:</strong> Ionic compound dissolved in water or molten</li>
<li><strong>Cathode (negative electrode):</strong> Cations (positive ions) migrate HERE; reduction (gain electrons)</li>
<li><strong>Anode (positive electrode):</strong> Anions (negative ions) migrate HERE; oxidation (lose electrons)</li>
</ul>
<p><strong>Electrolysis of molten ionic compound (e.g., molten NaCl):</strong></p>
<ul class="notes-list">
<li>Cathode: Na⁺ + e⁻ → Na (sodium deposited)</li>
<li>Anode: 2Cl⁻ → Cl₂ + 2e⁻ (chlorine gas produced)</li>
</ul>
<p><strong>Electrolysis of aqueous solutions [E]:</strong></p>
<ul class="notes-list">
<li>More complex: water also provides H⁺ and OH⁻ ions</li>
<li>Cathode: H⁺ discharged (as H₂) unless metal ion is below hydrogen in reactivity series</li>
<li>Anode: OH⁻ discharged (as O₂) unless halide ions present (then halogen produced)</li>
</ul>
<p><strong>Electrolysis of aqueous copper(II) sulfate:</strong></p>
<ul class="notes-list">
<li>Cathode: Cu²⁺ + 2e⁻ → Cu (copper deposited, blue colour fades)</li>
<li>Anode: 2H₂O → O₂ + 4H⁺ + 4e⁻ (oxygen gas)</li>
</ul>
<p><strong>Electroplating:</strong></p>
<ul class="notes-list">
<li>Object to plate = cathode; plating metal = anode</li>
<li>Solution = salt of plating metal</li>
<li>Used to: protect from corrosion, decorative purposes</li>
</ul>
<hr class="section-divider">
<h2 class="notes-h2">TOPIC 6: CHEMICAL ENERGETICS</h2>
<h3 class="notes-h3">6.1 Exothermic and Endothermic Reactions</h3>
<div class="table-wrap"><table class="notes-table">
<thead>
<tr><th>Type</th><th>Energy change</th><th>Temperature change</th><th>ΔH</th></tr>
</thead><tbody>
<tr><td>Exothermic</td><td>Energy released to surroundings</td><td>Increases</td><td>Negative (−)</td></tr>
<tr><td>Endothermic</td><td>Energy absorbed from surroundings</td><td>Decreases</td><td>Positive (+)</td></tr>
</tbody></table></div>
<p><strong>Exothermic examples:</strong> Combustion, neutralisation, oxidation, respiration, precipitation</p>
<p><strong>Endothermic examples:</strong> Thermal decomposition, dissolving ammonium nitrate in water, photosynthesis</p>
<h3 class="notes-h3">6.2 Bond Energies [E]</h3>
<ul class="notes-list">
<li>Energy is <strong>absorbed</strong> to break bonds</li>
<li>Energy is <strong>released</strong> when bonds form</li>
<li>ΔH = Energy absorbed (breaking bonds) − Energy released (forming bonds)</li>
<li>If negative: exothermic (more energy released than absorbed)</li>
</ul>
<hr class="section-divider">
<h2 class="notes-h2">TOPIC 7: CHEMICAL REACTIONS</h2>
<h3 class="notes-h3">7.1 Rate of Reaction</h3>
<p><strong>Definition:</strong> The change in concentration of reactant/product per unit time.</p>
<p><strong>Factors affecting rate:</strong></p>
<ul class="notes-list">
<li><strong>Concentration:</strong> More particles per unit volume → more frequent collisions → faster reaction</li>
<li><strong>Pressure (gases):</strong> Higher pressure → more frequent collisions → faster reaction</li>
<li><strong>Temperature:</strong> Higher temp → particles move faster → more frequent AND more energetic collisions → faster reaction</li>
<li><strong>Surface area:</strong> Smaller particle size → larger surface area → more collisions → faster reaction</li>
<li><strong>Catalyst:</strong> Provides alternative pathway with lower activation energy → faster reaction</li>
</ul>
<p><strong>Collision theory:</strong> Reactions occur when particles collide with sufficient energy (≥ activation energy) and correct orientation.</p>
<p><strong>Measuring rate:</strong></p>
<ul class="notes-list">
<li>Gas produced: collect in a gas syringe, measure volume at intervals</li>
<li>Precipitate: measure time for cross to disappear (turbidity)</li>
<li>Colour change: use colorimeter</li>
<li>Mass loss: weigh flask on balance (if gas escapes)</li>
</ul>
<p><strong>Catalysts:</strong> Speed up reactions without being used up. Are chemically unchanged at end.</p>
<ul class="notes-list">
<li>Examples: MnO₂ (decomposition of H₂O₂); Fe (Haber process); V₂O₅ (Contact process); Pt (car catalytic converters)</li>
</ul>
<h3 class="notes-h3">7.2 Reversible Reactions and Equilibrium [E]</h3>
<p>Some reactions are reversible (can go forward and backward): A + B ⇌ C + D</p>
<p><strong>Dynamic equilibrium:</strong> Forward and reverse rates are equal; concentrations remain constant. Reached in a closed system.</p>
<p><strong>Le Chatelier's Principle:</strong> If a change is made to a system at equilibrium, the equilibrium shifts to oppose that change.</p>
<div class="table-wrap"><table class="notes-table">
<thead>
<tr><th>Change</th><th>Shift</th></tr>
</thead><tbody>
<tr><td>Increase concentration of reactant</td><td>Shifts RIGHT (makes more products)</td></tr>
<tr><td>Increase concentration of product</td><td>Shifts LEFT</td></tr>
<tr><td>Increase temperature</td><td>Shifts in endothermic direction</td></tr>
<tr><td>Increase pressure (gases)</td><td>Shifts towards fewer moles of gas</td></tr>
<tr><td>Add catalyst</td><td>No shift — just reaches equilibrium faster</td></tr>
</tbody></table></div>
<hr class="section-divider">
<h2 class="notes-h2">TOPIC 8: ACIDS, BASES AND SALTS</h2>
<h3 class="notes-h3">8.1 Acids and Bases</h3>
<p><strong>Acids:</strong> Release H⁺ ions in aqueous solution (pH < 7)</p>
<p><strong>Bases:</strong> React with acids to form salt and water. Metal oxides and metal hydroxides are bases.</p>
<p><strong>Alkalis:</strong> Soluble bases; release OH⁻ ions in solution (pH > 7)</p>
<p><strong>pH scale:</strong> 0–7 (acid), 7 (neutral), 7–14 (alkali)</p>
<p><strong>Strong vs. weak acids:</strong></p>
<ul class="notes-list">
<li><strong>Strong acid:</strong> Fully ionises in water (e.g., HCl, H₂SO₄, HNO₃)</li>
<li><strong>Weak acid:</strong> Partially ionises (e.g., CH₃COOH, H₂CO₃)</li>
</ul>
<p><strong>Common acids:</strong></p>
<ul class="notes-list">
<li>Hydrochloric acid: HCl → H⁺ + Cl⁻</li>
<li>Sulfuric acid: H₂SO₄ → 2H⁺ + SO₄²⁻</li>
<li>Nitric acid: HNO₃ → H⁺ + NO₃⁻</li>
</ul>
<h3 class="notes-h3">8.2 Reactions of Acids</h3>
<p><strong>Acid + Metal → Salt + Hydrogen</strong></p>
<ul class="notes-list">
<li>Zn + H₂SO₄ → ZnSO₄ + H₂</li>
<li>(Metals above hydrogen in reactivity series react)</li>
</ul>
<p><strong>Acid + Metal oxide → Salt + Water</strong></p>
<ul class="notes-list">
<li>CuO + H₂SO₄ → CuSO₄ + H₂O</li>
</ul>
<p><strong>Acid + Metal hydroxide (neutralisation) → Salt + Water</strong></p>
<ul class="notes-list">
<li>NaOH + HCl → NaCl + H₂O</li>
<li>Ionic: H⁺ + OH⁻ → H₂O</li>
</ul>
<p><strong>Acid + Metal carbonate → Salt + Water + Carbon dioxide</strong></p>
<ul class="notes-list">
<li>CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂</li>
</ul>
<p><strong>Salt names:</strong></p>
<ul class="notes-list">
<li>HCl → chloride salts</li>
<li>H₂SO₄ → sulfate salts</li>
<li>HNO₃ → nitrate salts</li>
<li>H₂CO₃ → carbonate salts</li>
<li>H₃PO₄ → phosphate salts</li>
</ul>
<h3 class="notes-h3">8.3 Preparing Salts</h3>
<p><strong>Soluble salts:</strong></p>
<p><em>From acid + excess solid metal/oxide/carbonate:</em></p>
<ol class="notes-list">
<li>Add excess solid to warm acid</li>
<li>Filter off excess solid</li>
<li>Evaporate filtrate to get crystals</li>
</ol>
<p><em>Titration (acid + alkali):</em></p>
<ol class="notes-list">
<li>Use indicator to find exact volumes</li>
<li>Repeat without indicator using exact volumes</li>
<li>Evaporate to crystallise</li>
</ol>
<p><strong>Insoluble salts (precipitation):</strong></p>
<ul class="notes-list">
<li>Mix two aqueous solutions containing the required ions</li>
<li>Insoluble salt precipitates out</li>
<li>Filter, wash, dry</li>
</ul>
<p><strong>Example:</strong> Making BaSO₄ (insoluble)</p>
<ul class="notes-list">
<li>BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)</li>
<li>Filter off the white precipitate</li>
</ul>
<h3 class="notes-h3">8.4 Solubility Rules (for precipitation)</h3>
<div class="table-wrap"><table class="notes-table">
<thead>
<tr><th>Ion</th><th>Solubility</th></tr>
</thead><tbody>
<tr><td>All nitrates (NO₃⁻)</td><td>Soluble</td></tr>
<tr><td>All sodium (Na⁺), potassium (K⁺), ammonium (NH₄⁺) salts</td><td>Soluble</td></tr>
<tr><td>Most chlorides (Cl⁻)</td><td>Soluble EXCEPT AgCl, PbCl₂</td></tr>
<tr><td>Most sulfates (SO₄²⁻)</td><td>Soluble EXCEPT BaSO₄, PbSO₄, CaSO₄ (slightly)</td></tr>
<tr><td>Most carbonates (CO₃²⁻)</td><td>INSOLUBLE except Na, K, NH₄</td></tr>
<tr><td>Most hydroxides (OH⁻)</td><td>INSOLUBLE except Na, K, Ca (slightly)</td></tr>
</tbody></table></div>
<hr class="section-divider">
<h2 class="notes-h2">TOPIC 9: THE REACTIVITY SERIES</h2>
<h3 class="notes-h3">9.1 Reactivity Series (Most to Least Reactive)</h3>
<p><strong>Potassium > Sodium > Calcium > Magnesium > Aluminium > Carbon > Zinc > Iron > Tin > Lead > Hydrogen > Copper > Silver > Gold</strong></p>
<p>Memory aid: <strong>P</strong>lease <strong>S</strong>top <strong>C</strong>alling <strong>M</strong>e <strong>A</strong> <strong>C</strong>lown <strong>Z</strong>ach, <strong>I</strong> <strong>T</strong>hink <strong>L</strong>eading <strong>H</strong>eads <strong>C</strong>reates <strong>S</strong>tress for <strong>G</strong>oals</p>
<p><strong>Observations with water:</strong></p>
<ul class="notes-list">
<li>K, Na, Ca: react vigorously with cold water</li>
<li>Mg: reacts with steam; very slow with cold water</li>
<li>Zn, Fe: react with steam only</li>
<li>Cu, Ag, Au: do not react with water</li>
</ul>
<p><strong>Displacement reactions:</strong></p>
<ul class="notes-list">
<li>A more reactive metal displaces a less reactive metal from its salt solution</li>
<li>Fe + CuSO₄ → FeSO₄ + Cu (iron displaces copper; solution turns from blue to colourless; reddish copper deposits)</li>
<li>Cu + FeSO₄ → No reaction (copper less reactive than iron)</li>
</ul>
<h3 class="notes-h3">9.2 Extraction of Metals</h3>
<div class="table-wrap"><table class="notes-table">
<thead>
<tr><th>Metal</th><th>Method</th><th>Reason</th></tr>
</thead><tbody>
<tr><td>Potassium, Sodium, Calcium, Magnesium, Aluminium</td><td>Electrolysis of molten compounds</td><td>Too reactive to reduce with carbon</td></tr>
<tr><td>Zinc, Iron, Tin, Lead</td><td>Reduction with carbon (coke) in blast furnace or similar</td><td>Carbon can displace these metals</td></tr>
<tr><td>Copper, Silver, Gold</td><td>Found native OR heating/physical methods</td><td>Unreactive; naturally occurring</td></tr>
</tbody></table></div>
<p><strong>Extraction of iron (Blast Furnace):</strong></p>
<ul class="notes-list">
<li>Raw materials: iron ore (haematite, Fe₂O₃), coke (carbon), limestone (CaCO₃), hot air</li>
<li>Reactions:</li>
</ul>
<ol class="notes-list">
<li>C + O₂ → CO₂</li>
<li>CO₂ + C → 2CO (carbon monoxide = reducing agent)</li>
<li>Fe₂O₃ + 3CO → 2Fe + 3CO₂ (iron reduced)</li>
<li>CaCO₃ → CaO + CO₂ (limestone decomposes)</li>
<li>CaO + SiO₂ → CaSiO₃ (slag removes acidic impurities)</li>
</ol>
<p><strong>Extraction of aluminium (electrolysis):</strong></p>
<ul class="notes-list">
<li>Ore: bauxite (Al₂O₃)</li>
<li>Dissolved in molten cryolite (lowers melting point → saves energy)</li>
<li>Cathode: Al³⁺ + 3e⁻ → Al</li>
<li>Anode: 2O²⁻ → O₂ + 4e⁻</li>
<li>Carbon anodes burn away (react with O₂ produced) and need regular replacement</li>
</ul>
<h3 class="notes-h3">9.3 Rusting</h3>
<p><strong>Conditions:</strong> Iron rusts when BOTH water AND oxygen are present.</p>
<ul class="notes-list">
<li>Experiment proof: iron nails in boiled water (no O₂) → no rust; in dry air (no H₂O) → no rust; in air + water → rusts</li>
</ul>
<p><strong>Preventing rust:</strong></p>
<ul class="notes-list">
<li>Barrier methods: painting, oiling/greasing, plastic coating, tin-plating</li>
<li>Sacrificial protection: zinc coating (galvanising); magnesium blocks</li>
<li>Making stainless steel (alloy with chromium/nickel)</li>
</ul>
<hr class="section-divider">
<h2 class="notes-h2">TOPIC 10: ORGANIC CHEMISTRY</h2>
<h3 class="notes-h3">10.1 Alkanes (saturated hydrocarbons)</h3>
<p><strong>General formula:</strong> CₙH₂ₙ₊₂</p>
<div class="table-wrap"><table class="notes-table">
<thead>
<tr><th>Name</th><th>Formula</th><th>State at room temp</th></tr>
</thead><tbody>
<tr><td>Methane</td><td>CH₄</td><td>Gas</td></tr>
<tr><td>Ethane</td><td>C₂H₆</td><td>Gas</td></tr>
<tr><td>Propane</td><td>C₃H₈</td><td>Gas</td></tr>
<tr><td>Butane</td><td>C₄H₁₀</td><td>Gas</td></tr>
<tr><td>Pentane</td><td>C₅H₁₂</td><td>Liquid</td></tr>
</tbody></table></div>
<p><strong>Properties:</strong></p>
<ul class="notes-list">
<li>Non-polar; insoluble in water</li>
<li>Only C−C and C−H single bonds</li>
<li>Low reactivity (no double bonds)</li>
</ul>
<p><strong>Reactions:</strong></p>
<ul class="notes-list">
<li><strong>Combustion:</strong> Complete → CO₂ + H₂O; Incomplete → CO (+ soot)</li>
<li>CH₄ + 2O₂ → CO₂ + 2H₂O</li>
<li><strong>Substitution with halogens [E]:</strong> In UV light, H is replaced by halogen</li>
<li>CH₄ + Cl₂ → CH₃Cl + HCl (in UV light)</li>
</ul>
<h3 class="notes-h3">10.2 Alkenes (unsaturated hydrocarbons)</h3>
<p><strong>General formula:</strong> CₙH₂ₙ (contain one C=C double bond)</p>
<div class="table-wrap"><table class="notes-table">
<thead>
<tr><th>Name</th><th>Formula</th></tr>
</thead><tbody>
<tr><td>Ethene</td><td>C₂H₄</td></tr>
<tr><td>Propene</td><td>C₃H₆</td></tr>
<tr><td>Butene</td><td>C₄H₈</td></tr>
</tbody></table></div>
<p><strong>Properties:</strong></p>
<ul class="notes-list">
<li>More reactive than alkanes due to C=C double bond</li>
<li>Decolourise bromine water (test for C=C)</li>
</ul>
<p><strong>Reactions:</strong></p>
<ul class="notes-list">
<li><strong>Addition with hydrogen (hydrogenation):</strong> C₂H₄ + H₂ → C₂H₆ (catalyst: Ni, 150°C)</li>
<li><strong>Addition with bromine:</strong> C₂H₄ + Br₂ → C₂H₄Br₂ (1,2-dibromoethane); bromine water decolourised</li>
<li><strong>Addition with water (hydration):</strong> C₂H₄ + H₂O → C₂H₅OH (catalyst: H₃PO₄, 300°C) — industrial ethanol production</li>
<li><strong>Polymerisation:</strong> Many alkene monomers join → polymer</li>
</ul>
<h3 class="notes-h3">10.3 Alcohols</h3>
<p><strong>Functional group:</strong> −OH (hydroxyl group)</p>
<p><strong>General formula:</strong> CₙH₂ₙ₊₁OH</p>
<div class="table-wrap"><table class="notes-table">
<thead>
<tr><th>Name</th><th>Formula</th></tr>
</thead><tbody>
<tr><td>Methanol</td><td>CH₃OH</td></tr>
<tr><td>Ethanol</td><td>C₂H₅OH</td></tr>
<tr><td>Propanol</td><td>C₃H₇OH</td></tr>
</tbody></table></div>
<p><strong>Production of ethanol:</strong></p>
<ol class="notes-list">
<li><strong>Fermentation:</strong> glucose → ethanol + CO₂ (yeast, anaerobic, ~30°C)</li>
</ol>
<ul class="notes-list">
<li>C₆H₁₂O₆ → 2C₂H₅OH + 2CO₂</li>
</ul>
<ol class="notes-list">
<li><strong>Hydration of ethene</strong> (industrial, faster, continuous process)</li>
</ol>
<p><strong>Reactions of ethanol:</strong></p>
<ul class="notes-list">
<li>Combustion: C₂H₅OH + 3O₂ → 2CO₂ + 3H₂O</li>
<li>Oxidation: ethanol → ethanoic acid (by bacteria or oxidising agents)</li>
</ul>
<h3 class="notes-h3">10.4 Carboxylic Acids</h3>
<p><strong>Functional group:</strong> −COOH (carboxyl)</p>
<div class="table-wrap"><table class="notes-table">
<thead>
<tr><th>Name</th><th>Formula</th></tr>
</thead><tbody>
<tr><td>Methanoic acid</td><td>HCOOH</td></tr>
<tr><td>Ethanoic acid</td><td>CH₃COOH</td></tr>
</tbody></table></div>
<ul class="notes-list">
<li>Weak acids (partially ionise)</li>
<li>React with alkalis, carbonates (like any acid)</li>
<li>Ethanoic acid = acetic acid (vinegar)</li>
</ul>
<h3 class="notes-h3">10.5 Polymers [E]</h3>
<p><strong>Addition polymers:</strong> Alkene monomers join (no other product)</p>
<ul class="notes-list">
<li>Ethene → poly(ethene): nCH₂=CH₂ → (−CH₂−CH₂−)ₙ</li>
<li>Propene → poly(propene)</li>
<li>Chloroethene → PVC</li>
<li>Tetrafluoroethene → PTFE (Teflon)</li>
</ul>
<p><strong>Condensation polymers:</strong> Monomers join with loss of small molecule (usually H₂O)</p>
<ul class="notes-list">
<li><strong>Polyesters:</strong> Diol + dicarboxylic acid → polyester + H₂O (e.g., Terylene)</li>
<li><strong>Nylon (polyamide):</strong> Diamine + dicarboxylic acid → nylon + H₂O</li>
</ul>
<p><strong>Problems with plastics:</strong></p>
<ul class="notes-list">
<li>Non-biodegradable (persist in environment)</li>
<li>Toxic gases when burned</li>
<li>Solutions: recycling, biodegradable plastics, reducing use</li>
</ul>
<h3 class="notes-h3">10.6 Crude Oil and Fractions</h3>
<p><strong>Fractional distillation of crude oil:</strong></p>
<ul class="notes-list">
<li>Crude oil is a mixture of hydrocarbons</li>
<li>Separated by boiling point in fractionating column</li>
</ul>
<div class="table-wrap"><table class="notes-table">
<thead>
<tr><th>Fraction</th><th>Carbon atoms</th><th>Boiling point range</th><th>Uses</th></tr>
</thead><tbody>
<tr><td>Refinery gases</td><td>C₁–C₄</td><td>Below 25°C</td><td>Fuel gas, LPG</td></tr>
<tr><td>Petrol (gasoline)</td><td>C₅–C₁₀</td><td>40–75°C</td><td>Car fuel</td></tr>
<tr><td>Naphtha</td><td>C₆–C₁₀</td><td>75–120°C</td><td>Chemicals feedstock</td></tr>
<tr><td>Kerosene</td><td>C₁₀–C₁₆</td><td>120–240°C</td><td>Jet fuel</td></tr>
<tr><td>Diesel oil</td><td>C₁₄–C₁₉</td><td>240–350°C</td><td>Diesel engines</td></tr>
<tr><td>Fuel oil</td><td>C₁₉–C₂₅</td><td>350–400°C</td><td>Ships, power stations</td></tr>
<tr><td>Bitumen</td><td>C₂₅+</td><td>400°C+</td><td>Roads, roofing</td></tr>
</tbody></table></div>
<p><strong>Cracking [E]:</strong></p>
<ul class="notes-list">
<li>Long-chain alkanes broken into shorter, more useful alkenes and alkanes</li>
<li>Thermal cracking: high temperature and pressure</li>
<li>Catalytic cracking: zeolite catalyst, moderate temperature</li>
<li>Example: C₁₀H₂₂ → C₅H₁₂ + C₄H₈ + CH₄</li>
</ul>
<hr class="section-divider">
<h2 class="notes-h2">TOPIC 11: EXPERIMENTAL TECHNIQUES</h2>
<h3 class="notes-h3">11.1 Separation Techniques</h3>
<div class="table-wrap"><table class="notes-table">
<thead>
<tr><th>Technique</th><th>Separates</th><th>Method</th></tr>
</thead><tbody>
<tr><td>Filtration</td><td>Insoluble solid from liquid</td><td>Filter paper; solid stays on paper</td></tr>
<tr><td>Evaporation</td><td>Dissolved solid from liquid</td><td>Heat solution; liquid evaporates</td></tr>
<tr><td>Crystallisation</td><td>Dissolved solid from solution</td><td>Heat to concentrate, then cool slowly</td></tr>
<tr><td>Distillation</td><td>Liquid mixture by boiling point</td><td>Heat; collect vapour at different temperatures</td></tr>
<tr><td>Fractional distillation</td><td>Liquids with close boiling points</td><td>Column with fractionating column</td></tr>
<tr><td>Chromatography</td><td>Dissolved substances by solubility</td><td>Different Rf values on paper/plate</td></tr>
</tbody></table></div>
<h3 class="notes-h3">11.2 Chromatography</h3>
<p><strong>Rf value = distance moved by spot / distance moved by solvent front</strong></p>
<ul class="notes-list">
<li>Substances with higher solubility in solvent travel further (higher Rf)</li>
<li>Rf values are constant for a given solvent and stationary phase</li>
<li>Used to identify unknown substances by comparing Rf values to known standards</li>
</ul>
<h3 class="notes-h3">11.3 Tests for Gases</h3>
<div class="table-wrap"><table class="notes-table">
<thead>
<tr><th>Gas</th><th>Test</th><th>Positive result</th></tr>
</thead><tbody>
<tr><td>Hydrogen</td><td>Lit splint</td><td>Squeaky "pop"</td></tr>
<tr><td>Oxygen</td><td>Glowing splint</td><td>Splint relights</td></tr>
<tr><td>Carbon dioxide</td><td>Limewater</td><td>Turns cloudy/milky</td></tr>
<tr><td>Chlorine</td><td>Damp litmus paper</td><td>Bleaches/turns white then red</td></tr>
<tr><td>Ammonia</td><td>Damp red litmus</td><td>Turns blue</td></tr>
</tbody></table></div>
<h3 class="notes-h3">11.4 Tests for Ions [E]</h3>
<p><strong>Flame test for cations:</strong></p>
<div class="table-wrap"><table class="notes-table">
<thead>
<tr><th>Ion</th><th>Flame colour</th></tr>
</thead><tbody>
<tr><td>Li⁺</td><td>Crimson/red</td></tr>
<tr><td>Na⁺</td><td>Yellow</td></tr>
<tr><td>K⁺</td><td>Lilac/violet</td></tr>
<tr><td>Ca²⁺</td><td>Brick red/orange</td></tr>
<tr><td>Cu²⁺</td><td>Blue-green</td></tr>
</tbody></table></div>
<p><strong>Aqueous tests for cations (add NaOH solution):</strong></p>
<div class="table-wrap"><table class="notes-table">
<thead>
<tr><th>Ion</th><th>Observation</th><th>Further: add excess NaOH</th></tr>
</thead><tbody>
<tr><td>Fe²⁺</td><td>Green precipitate</td><td>No change</td></tr>
<tr><td>Fe³⁺</td><td>Orange/brown precipitate</td><td>No change</td></tr>
<tr><td>Cu²⁺</td><td>Blue precipitate</td><td>No change</td></tr>
<tr><td>NH₄⁺</td><td>No precipitate</td><td>Ammonia gas released (pungent smell; turns damp red litmus blue)</td></tr>
<tr><td>Al³⁺, Zn²⁺</td><td>White precipitate</td><td>Dissolves (amphoteric)</td></tr>
</tbody></table></div>
<p><strong>Tests for anions:</strong></p>
<div class="table-wrap"><table class="notes-table">
<thead>
<tr><th>Ion</th><th>Test</th><th>Positive result</th></tr>
</thead><tbody>
<tr><td>Cl⁻</td><td>Add dilute HNO₃, then AgNO₃</td><td>White precipitate (AgCl)</td></tr>
<tr><td>Br⁻</td><td>Add dilute HNO₃, then AgNO₃</td><td>Cream precipitate (AgBr)</td></tr>
<tr><td>I⁻</td><td>Add dilute HNO₃, then AgNO₃</td><td>Yellow precipitate (AgI)</td></tr>
<tr><td>SO₄²⁻</td><td>Add dilute HCl, then BaCl₂</td><td>White precipitate (BaSO₄)</td></tr>
<tr><td>CO₃²⁻</td><td>Add dilute HCl</td><td>Effervescence (CO₂ produced)</td></tr>
</tbody></table></div>
<hr class="section-divider">
<h2 class="notes-h2">EXAM TECHNIQUE GUIDE</h2>
<h3 class="notes-h3">Common Mistakes to Avoid</h3>
<ol class="notes-list">
<li><strong>Electrolysis:</strong> Remember cathode = negative (cations go there); anode = positive (anions go there). Mnemonics: "AN OX RED CAT" (Anode = Oxidation; Cathode = Reduction)</li>
<li><strong>Rusting:</strong> BOTH water AND oxygen needed — not just one</li>
<li><strong>Salt names:</strong> HCl → chloride (NOT hydrochlorate); H₂SO₄ → sulfate; HNO₃ → nitrate</li>
<li><strong>Balancing equations:</strong> Count EVERY atom; balance one element at a time</li>
<li><strong>Mole calculations:</strong> Always write n = m/Mᵣ formula first; show working</li>
<li><strong>Alkenes vs. alkanes:</strong> Alkenes have C=C (more reactive; decolourise Br₂ water); alkanes do not</li>
<li><strong>Ionic equations:</strong> Only show species that change; spectator ions are NOT included</li>
<li><strong>State symbols:</strong> Always include (s), (l), (g), (aq) in full chemical equations</li>
</ol>
<h3 class="notes-h3">Mark-Winning Phrases</h3>
<ul class="notes-list">
<li>For rate of reaction: "…increasing temperature causes particles to move faster, so there are more frequent collisions AND more particles have energy ≥ activation energy"</li>
<li>For electrolysis cathode: "…positive ions gain electrons (are reduced) at the cathode"</li>
<li>For extraction: "…aluminium is extracted by electrolysis because it is too reactive to be reduced by carbon"</li>
<li>For rusting prevention: "…zinc coating provides barrier protection AND acts as a sacrificial metal (more reactive than iron, so it corrodes preferentially)"</li>
<li>For addition polymerisation: "…the C=C double bond in ethene opens up and monomers join together forming long polymer chains with no other products"</li>
</ul>
<hr class="section-divider">
<h2 class="notes-h2">KEY FORMULAE</h2>
<pre class="code-block"><code>
n (moles) = mass (g) ÷ Mᵣ
n = volume (dm³) × concentration (mol/dm³)
n (gas) = volume (dm³) ÷ 24 [at RTP]
Concentration = n ÷ volume (dm³)
% yield = (actual yield ÷ theoretical yield) × 100
% purity = (mass of pure substance ÷ total mass) × 100
Rf = distance moved by spot ÷ distance moved by solvent
</code></pre>
<hr class="section-divider">
<p><em>These notes cover the complete Cambridge IGCSE Chemistry (0620) syllabus.</em></p>
<p><em>For official past papers and mark schemes: cambridgeinternational.org</em></p>
<p><em>For full IGCSE Chemistry revision bundle: cbcedukenya.com/igcse-chemistry</em></p>
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