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Chemistry
Study Notes
All notes follow the official WAEC and JAMB approved syllabus. Three sections: Physical Chemistry, Inorganic Chemistry and Organic Chemistry. Study a topic first, then take the practice quiz to find your weak areas.
Separation of Mixtures
Distillation, chromatography, filtration, sublimation
Atomic Structure & Bonding
Electron config, periodic table, ionic, covalent, metallic bonds
Kinetic Theory & Gas Laws
Boyle's, Charles's, Graham's, ideal gas equation
Stoichiometry & Mole Concept
Molar mass, Avogadro's number, empirical formula, titration
Solutions & Solubility
Concentration, solubility curves, colligative properties
Electrolysis
Faraday's laws, electrolytic cells, electroplating
Energy Changes
Exothermic, endothermic, enthalpy, Hess's law
Rates & Equilibrium
Factors affecting rate, Le Chatelier's principle
Oxidation & Reduction
Oxidation states, redox reactions, half equations
Periodic Table
Groups, periods, trends in properties, periodicity
Acids, Bases & Salts
pH, neutralisation, preparation of salts
Air, Water & Pollution
Composition of air, water purification, pollutants
Metals & Their Compounds
Extraction, alloys, Group I & II, transition metals
Non-Metals & Compounds
Hydrogen, oxygen, nitrogen, halogens, sulphur
Intro to Organic Chemistry
Carbon bonding, homologous series, functional groups, isomerism
Hydrocarbons
Alkanes, alkenes, alkynes — naming, reactions, petroleum
Functional Group Compounds
Alkanols, alkanoic acids, esters, amines, aldehydes, ketones
Chemistry & Industry
Haber process, Contact process, fertilisers, polymers, soap
Separation of Mixtures & Purification
A pure substance has a sharp melting point and boiling point. An impure substance has a depressed melting point and elevated boiling point. Mixtures can be separated by physical methods since their components retain their properties.
| Method | Principle | Used For | Example |
|---|---|---|---|
| Filtration | Separates an insoluble solid from a liquid | Solid + liquid mixture | Sand from water |
| Evaporation | Evaporates solvent to leave dissolved solid | Solution where solid is required | Salt from salt water |
| Simple distillation | Boiling and condensing; separates components by boiling point | Solution of dissolved solids; two liquids with very different b.p. | Water from salt solution; ethanol from water |
| Fractional distillation | Separation of liquids with close boiling points using a fractionating column | Miscible liquids close in b.p. | Crude oil refining; air separation; ethanol/water mixtures |
| Crystallisation | Cooling a saturated hot solution to form crystals | Soluble solid from solution (more pure than evaporation) | Copper(II) sulphate crystals from solution |
| Sublimation | Solid converts directly to gas (skips liquid phase) | Separating a sublimable solid from non-sublimable solids | Iodine from sand; ammonium chloride from salt |
| Magnetisation | Uses a magnet to attract magnetic materials | Magnetic solid from non-magnetic solid | Iron filings from sulphur |
| Decantation | Carefully pouring off liquid from settled solid | Coarse solid that has settled from liquid | Sand and water after settling |
| Paper chromatography | Different components travel different distances up paper depending on solubility | Separating coloured pigments, inks, dyes | Separating inks in a pen |
| Solvent extraction | Using a selective solvent to dissolve one component | Separating based on solubility difference | Iodine in organic solvent from water |
WAEC and JAMB test sublimation and chromatography frequently. Know: sublimation separates iodine from sand or salt. The Rf value in chromatography = distance moved by spot ÷ distance moved by solvent. Pure substances have one spot; mixtures have multiple spots.
Atomic Structure & Chemical Bonding
Atom: smallest particle of an element that retains chemical properties. Consists of a nucleus (protons + neutrons) surrounded by electrons in shells (energy levels).
| Particle | Charge | Relative Mass | Location |
|---|---|---|---|
| Proton | +1 | 1 | Nucleus |
| Neutron | 0 | 1 | Nucleus |
| Electron | −1 | 1/1836 (negligible) | Shells/orbitals around nucleus |
- Atomic number (Z) = number of protons (= number of electrons in a neutral atom)
- Mass number (A) = protons + neutrons
- Neutrons = mass number − atomic number
- Isotopes: atoms of the same element with same atomic number but different mass numbers (different number of neutrons). E.g. ¹²C and ¹⁴C
Shells hold: shell 1 = max 2 electrons · shell 2 = max 8 · shell 3 = max 8 (for elements 1–20)
Na (11): 2, 8, 1 · Cl (17): 2, 8, 7 · Ca (20): 2, 8, 8, 2
The valence electrons (outer shell) determine chemical properties and bonding.
| Bond Type | How it forms | Properties of compounds | Example |
|---|---|---|---|
| Ionic (electrovalent) | Transfer of electrons from metal to non-metal. Both achieve noble gas configuration. | High m.p./b.p., conducts when molten or dissolved, brittle, usually soluble in water | NaCl, MgO, CaCl₂ |
| Covalent | Sharing of electrons between non-metals | Usually low m.p./b.p., doesn't conduct electricity (except graphite), often volatile | H₂O, CO₂, CH₄, HCl |
| Co-ordinate (dative) covalent | Both electrons in the shared pair come from ONE atom (the donor) | Similar to covalent | NH₄⁺ (ammonium ion), H₃O⁺ |
| Metallic | Positive metal ions in a "sea" of delocalised electrons | Good conductor of heat and electricity, malleable, ductile, high m.p. | All metals: Fe, Cu, Na, Al |
| Hydrogen bonding | Electrostatic attraction between δ+ hydrogen and electronegative atom (N, O, F) on another molecule | Higher than expected b.p./m.p., surface tension in water | H₂O, HF, NH₃, DNA structure |
Ionic vs Covalent: Ionic = metal + non-metal. Covalent = non-metal + non-metal. This rule covers ~95% of WAEC/JAMB questions on bonding type. Water's anomalously high boiling point (100°C vs expected ~−80°C) is due to hydrogen bonding.
Kinetic Theory & Gas Laws
All matter is made of particles in constant random motion. The higher the temperature, the faster the particles move. Evidence: Brownian motion (random zig-zag movement of small particles) and diffusion (spontaneous mixing of particles).
Graham's Law of Diffusion: The rate of diffusion of a gas is inversely proportional to the square root of its molar mass. Lighter gases diffuse faster.
NH₃ (molar mass 17) diffuses faster than HCl (molar mass 36.5). That's why in the white ring experiment, the ring forms closer to HCl end.
| Law | Statement | Formula | Condition |
|---|---|---|---|
| Boyle's Law | At constant temperature, pressure and volume of a fixed mass of gas are inversely proportional | P₁V₁ = P₂V₂ | Constant temperature (isothermal) |
| Charles's Law | At constant pressure, volume of a fixed mass of gas is directly proportional to absolute temperature | V₁/T₁ = V₂/T₂ | Constant pressure (isobaric) |
| Pressure Law (Gay-Lussac) | At constant volume, pressure is directly proportional to absolute temperature | P₁/T₁ = P₂/T₂ | Constant volume (isochoric) |
| General Gas Equation | Combines Boyle's and Charles's laws | P₁V₁/T₁ = P₂V₂/T₂ | Fixed mass of gas |
| Ideal Gas Equation | Relates all gas properties | PV = nRT | n = moles, R = 8.314 J/mol·K |
Always convert temperature to Kelvin (K): K = °C + 273. If you use °C in gas law equations, you WILL get the wrong answer. STP = 0°C (273 K) and 1 atm (101.325 kPa). Molar volume of gas at STP = 22.4 dm³/mol.
Chemical Combination, Mole Concept & Stoichiometry
- Law of Conservation of Mass: Matter is neither created nor destroyed in a chemical reaction. Total mass of reactants = total mass of products.
- Law of Definite Proportions (Constant Composition): A compound always contains the same elements in the same ratio by mass. Water is always H:O = 1:8 by mass.
- Law of Multiple Proportions: When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in simple whole-number ratios. CO and CO₂ example.
- Avogadro's Law: Equal volumes of all gases at the same temperature and pressure contain equal numbers of molecules.
- 1 mole of any substance contains 6.02 × 10²³ particles (Avogadro's constant, Nₐ).
- Molar mass = mass of 1 mole of a substance in grams (= Ar or Mr in g/mol).
- Moles = mass (g) ÷ molar mass (g/mol)
- Moles of gas = volume (dm³) ÷ 22.4 (at STP)
- Concentration (mol/dm³) = moles ÷ volume (dm³)
Step 1: Convert percentages to grams. Step 2: Divide each by atomic mass to get moles. Step 3: Divide all moles by the smallest. Step 4: Round to whole numbers → Empirical formula. Step 5: If molecular mass given: n = Mr ÷ empirical formula mass → multiply empirical by n.
Titration calculation: C₁V₁/n₁ = C₂V₂/n₂ where n₁ and n₂ are the mole ratios from the balanced equation. This formula applies to acid-base titrations and is tested in virtually every WAEC Paper 2.
Solutions & Solubility
- Solution: A homogeneous mixture of solute dissolved in solvent.
- Solubility: Mass of solute that dissolves in 100 g of solvent at a specified temperature to form a saturated solution.
- Saturated solution: Cannot dissolve more solute at that temperature.
- Supersaturated: Contains more dissolved solute than normal saturation allows (unstable).
- Miscible: Two liquids that mix completely (e.g. ethanol + water).
- Immiscible: Two liquids that do not mix (e.g. oil + water).
Most solids: solubility increases with temperature (e.g. KNO₃, NaCl slightly).
Gases: solubility decreases with temperature — why fizzy drinks go flat when warm.
From a solubility curve: mass of crystals deposited = mass dissolved at T₁ − mass dissolved at T₂.
Colligative properties (WAEC Paper 2): Boiling point elevation, freezing point depression, osmotic pressure. These properties depend on the NUMBER of particles in solution, not their identity. Adding solute raises boiling point and lowers freezing/melting point — this is why salt is added to icy roads.
Electrolysis
Electrolysis is the decomposition of an ionic compound in molten or aqueous state by passing electricity through it. Requires an electrolyte (ionic conductor), a cathode (negative electrode) and an anode (positive electrode).
- Cathode (−): Attracts cations (positive ions). Reduction occurs here (gain of electrons).
- Anode (+): Attracts anions (negative ions). Oxidation occurs here (loss of electrons).
Oxidation Is Loss (of electrons) → occurs at ANODE
Reduction Is Gain (of electrons) → occurs at CATHODE
| Electrolyte | At Cathode (−) | At Anode (+) |
|---|---|---|
| Molten NaCl | Na metal deposited | Cl₂ gas evolved |
| Dilute H₂SO₄ | H₂ gas evolved | O₂ gas evolved |
| CuSO₄ (inert anode) | Cu deposited | O₂ evolved |
| CuSO₄ (Cu anode) | Cu deposited | Cu dissolves (anode loses mass) |
First Law: Mass of substance deposited is directly proportional to the quantity of electricity (charge) passed. Q = It (charge = current × time).
Second Law: Masses of different substances deposited by the same charge are proportional to their chemical equivalents (molar mass ÷ valency).
Mass deposited = (M × I × t) ÷ (n × F)
M = molar mass, I = current (A), t = time (s), n = number of electrons transferred, F = Faraday's constant = 96,500 C/mol
Applications of electrolysis: Electroplating (coating metal object with another metal), extraction of reactive metals (sodium by Downs process, aluminium by Hall-Héroult process), purification of copper, manufacture of chlorine and NaOH (chlor-alkali industry).
Energy Changes in Chemical Reactions
| Feature | Exothermic | Endothermic |
|---|---|---|
| Energy change | Releases energy (heat) to surroundings | Absorbs energy from surroundings |
| Temperature of surroundings | Increases (feels hot) | Decreases (feels cold) |
| Enthalpy change (ΔH) | Negative (ΔH < 0) | Positive (ΔH > 0) |
| Energy profile | Products lower than reactants | Products higher than reactants |
| Examples | Combustion, neutralisation, respiration, rusting | Photosynthesis, thermal decomposition, dissolving NH₄NO₃ |
The enthalpy change of a reaction is the same regardless of the route taken, provided initial and final conditions are the same. This allows us to calculate enthalpy changes that cannot be measured directly by using a cycle of known reactions.
Bond energy: Breaking bonds requires energy (endothermic). Forming bonds releases energy (exothermic). ΔH = Energy to break bonds − Energy released forming bonds. If negative → exothermic reaction overall.
Rates of Chemical Reaction & Chemical Equilibrium
| Factor | Effect on Rate | Explanation |
|---|---|---|
| Temperature | Increases rate | More particles have activation energy; more frequent effective collisions |
| Concentration (solution) | Increases rate | More particles per unit volume; more frequent collisions |
| Pressure (gases) | Increases rate | Equivalent to increased concentration of gas particles |
| Surface area | Increases rate | More surface exposed for collisions (powder reacts faster than lumps) |
| Catalyst | Increases rate | Provides alternative reaction pathway with lower activation energy; NOT consumed |
| Light | Can increase rate | Photochemical reactions (e.g. photosynthesis, photography) |
Collision Theory: Reactions occur when particles collide with sufficient energy (≥ activation energy) AND correct orientation. A catalyst LOWERS activation energy; it does NOT increase the frequency of collisions or change enthalpy.
A reversible reaction reaches dynamic equilibrium when the forward and reverse reaction rates are equal, and concentrations of reactants and products remain constant.
Le Chatelier's Principle: If a stress (change) is applied to a system at equilibrium, the system will shift to oppose that change.
| Change Applied | Direction Equilibrium Shifts |
|---|---|
| Increase concentration of reactants | Forward (→) to use up the reactants |
| Increase pressure (gas reactions) | Towards side with FEWER moles of gas |
| Increase temperature | Endothermic direction (absorbs heat to reduce temperature) |
| Add a catalyst | No shift — equilibrium reached faster only |
High pressure → shifts right (4 moles → 2 moles). Low temperature → more NH₃ (exothermic reaction). But in practice, 450°C is used as a compromise — low temperature gives high yield but unacceptably slow rate. Iron catalyst speeds up the reaction.
Oxidation & Reduction (Redox)
| Definition | Oxidation | Reduction |
|---|---|---|
| Classical (oxygen) | Gain of oxygen OR loss of hydrogen | Loss of oxygen OR gain of hydrogen |
| Electron transfer | Loss of electrons (LEO) | Gain of electrons (GER) |
| Oxidation state | Increase in oxidation state | Decrease in oxidation state |
Oxidation Is Loss of electrons | Reduction Is Gain of electrons
The substance that loses electrons is oxidised (it is the reducing agent).
The substance that gains electrons is reduced (it is the oxidising agent).
Oxidation state rules: Oxygen = −2 (except in peroxides = −1). Hydrogen = +1 (except in metal hydrides = −1). Sum of oxidation states in a compound = 0. In an ion = the charge of the ion. Use these to assign oxidation states to all atoms in any formula.
Periodic Table & Periodicity
The modern periodic table arranges elements in order of increasing atomic number. Elements in the same group have the same number of valence electrons and similar chemical properties. Elements in the same period have the same number of electron shells.
| Property | Across Period (left → right) | Down Group (top → bottom) |
|---|---|---|
| Atomic radius | Decreases (more protons pull electrons closer) | Increases (more electron shells) |
| Ionisation energy | Increases (harder to remove e⁻ from smaller atom) | Decreases (outer electrons further from nucleus) |
| Electronegativity | Increases (F is most electronegative) | Decreases |
| Electron affinity | Increases | Decreases |
| Metallic character | Decreases (metals on left, non-metals on right) | Increases |
Group I (Alkali metals): Li, Na, K — very reactive, react with water, form +1 ions, hydroxides and oxides.
Group II (Alkaline earth metals): Mg, Ca — less reactive than Group I, form +2 ions.
Group VII (Halogens): F, Cl, Br, I — reactive non-metals, form −1 ions, reactivity decreases down group.
Group 0/VIII (Noble gases): He, Ne, Ar — unreactive, full outer shells, used in lighting and balloons.
Acids, Bases & Salts
| Theory | Acid | Base |
|---|---|---|
| Arrhenius | Produces H⁺ ions in water | Produces OH⁻ ions in water |
| Brønsted-Lowry | Proton (H⁺) donor | Proton (H⁺) acceptor |
| Lewis | Electron pair acceptor | Electron pair donor |
pH 0–6 = acidic. pH 7 = neutral. pH 8–14 = alkaline/basic.
Strong acids: HCl, H₂SO₄, HNO₃ — fully ionise in water.
Weak acids: CH₃COOH (ethanoic acid), H₂CO₃ — partially ionise.
| Salt Type | Method of Preparation | Example |
|---|---|---|
| Soluble salt (from reactive metal/carbonate) | Metal or metal carbonate + acid; filter excess solid; evaporate | ZnSO₄: Zn + H₂SO₄; MgCl₂: MgCO₃ + HCl |
| Soluble salt (from alkali + acid) | Titration — exact volumes mixed and evaporated | NaCl: NaOH + HCl |
| Insoluble salt | Precipitation — mix two soluble solutions | BaSO₄: BaCl₂ + H₂SO₄; AgCl: AgNO₃ + HCl |
Hydrolysis of salts: Salts from strong acid + strong base = neutral. Strong acid + weak base = acidic (e.g. NH₄Cl). Weak acid + strong base = alkaline (e.g. Na₂CO₃). This determines whether a salt solution is acidic, neutral or alkaline.
Air, Water & Environmental Pollution
| Component | Approximate % by volume |
|---|---|
| Nitrogen (N₂) | 78% |
| Oxygen (O₂) | 21% |
| Argon and other noble gases | ~0.93% |
| Carbon dioxide (CO₂) | ~0.04% |
| Water vapour, dust, etc. | Variable |
Air pollutants and their effects: CO (carbon monoxide) — colourless, odourless, toxic — binds haemoglobin. SO₂ — causes acid rain, respiratory issues. NO₂ — acid rain, smog. CO₂ — greenhouse effect/global warming. Lead compounds — neurotoxic. CFCs — ozone layer depletion.
- Hard water: Contains dissolved Ca²⁺ and Mg²⁺ ions (from limestone/gypsum areas). Does not lather easily with soap.
- Temporary hardness: Caused by Ca(HCO₃)₂. Removed by boiling.
- Permanent hardness: Caused by CaSO₄, MgSO₄. Not removed by boiling. Removed by adding washing soda (Na₂CO₃) or by ion exchange.
- Water purification: Sedimentation → coagulation (alum added) → filtration through sand bed → chlorination (kills bacteria) → storage.
Metals & Their Compounds
K > Na > Ca > Mg > Al > Zn > Fe > Sn > Pb > H > Cu > Ag > Au
Memory: King Napoleon Came Marching Along Zulu Fields Singing Powerful Hymns, Cuute Silver Auburn
- Metals above hydrogen react with dilute acids to produce hydrogen gas.
- More reactive metals displace less reactive metals from salt solutions (displacement reaction).
- Very reactive metals (K, Na, Ca) react vigorously with cold water. Mg reacts with steam. Al, Zn, Fe react with steam. Cu, Ag, Au do not react with water.
| Metal | Method of Extraction | Ore |
|---|---|---|
| Sodium (Na) | Electrolysis of molten NaCl (Downs process) | Rock salt (NaCl) |
| Aluminium (Al) | Electrolysis of molten Al₂O₃ (Hall-Héroult process) | Bauxite (Al₂O₃·2H₂O) |
| Iron (Fe) | Reduction by coke (carbon) in blast furnace | Haematite (Fe₂O₃), Magnetite (Fe₃O₄) |
| Copper (Cu) | Smelting then electrolytic purification | Chalcopyrite (CuFeS₂) |
| Zinc (Zn) | Reduction by carbon | Zinc blende (ZnS) |
Rule: The MORE reactive the metal, the MORE energy (electrolysis) is needed to extract it. Less reactive metals are extracted by carbon reduction. The LEAST reactive (Ag, Au) are found native (as free elements in the Earth's crust).
Non-Metals & Their Compounds
| Non-Metal | Key Facts | Important Compounds |
|---|---|---|
| Hydrogen (H₂) | Lightest element. Burns in air to form water. Reduces metal oxides. Used in Haber process, hydrogenation of oils. | Water (H₂O), HCl, H₂SO₄, NH₃ |
| Oxygen (O₂) | Makes up 21% of air. Supports combustion. Made by electrolysis of water or decomposition of H₂O₂. | Oxides, water, CO₂ |
| Nitrogen (N₂) | Makes up 78% of air. Very unreactive due to N≡N triple bond. Used in Haber process to make ammonia. | NH₃, NO₂, HNO₃, N₂O |
| Chlorine (Cl₂) | Yellow-green gas. Toxic. Strong oxidising agent. Bleaches moist litmus. Made in chlor-alkali industry. | HCl, NaCl, bleaching powder, PVC |
| Sulphur (S) | Yellow solid. Burns to SO₂. Used in Contact process (H₂SO₄ manufacture) and vulcanisation of rubber. | SO₂, SO₃, H₂SO₄, H₂S |
| Carbon (C) | Allotropes: diamond (hardest), graphite (conductor), fullerene. CO is toxic; CO₂ is greenhouse gas. | CO, CO₂, carbonates, organic compounds |
Ammonia (NH₃): Made by Haber process (N₂ + 3H₂ ⇌ 2NH₃). Properties: colourless gas, pungent smell, lighter than air, turns moist red litmus blue, alkaline in water. Used in fertilisers (NH₄NO₃, (NH₄)₂SO₄, urea), explosives, cleaning agents.
Introduction to Organic Chemistry
Organic chemistry is the study of compounds containing carbon (with few exceptions like CO, CO₂, carbonates). Carbon forms 4 bonds and can bond to itself in chains and rings — making millions of possible compounds.
Homologous series: A family of organic compounds with the same general formula and functional group, differing only by CH₂. They show a gradual change in physical properties and similar chemical properties.
| Functional Group | Compound Type | General Formula | Example |
|---|---|---|---|
| —OH (hydroxyl) | Alkanols (alcohols) | CₙH₂ₙ₊₂O | CH₃OH (methanol), C₂H₅OH (ethanol) |
| —COOH (carboxyl) | Alkanoic acids (carboxylic acids) | CₙH₂ₙO₂ | CH₃COOH (ethanoic acid/acetic acid) |
| —COO— (ester linkage) | Alkanoates (esters) | CₙH₂ₙO₂ | CH₃COOC₂H₅ (ethyl ethanoate) |
| —NH₂ (amino) | Amines | CₙH₂ₙ₊₃N | CH₃NH₂ (methylamine) |
| —CHO (aldehyde) | Alkanals (aldehydes) | CₙH₂ₙO | HCHO (methanal/formaldehyde) |
| >C=O (ketone) | Alkanones (ketones) | CₙH₂ₙO | CH₃COCH₃ (propanone/acetone) |
Isomers: Compounds with the same molecular formula but different structural arrangements. Types: structural isomers (different carbon skeleton or functional group position); geometric isomers (cis/trans — in alkenes). Isomers have same molecular formula but different properties.
Hydrocarbons: Alkanes, Alkenes & Alkynes
| Feature | Alkanes | Alkenes | Alkynes |
|---|---|---|---|
| General formula | CₙH₂ₙ₊₂ | CₙH₂ₙ | CₙH₂ₙ₋₂ |
| Bond type | C—C single bonds only (saturated) | C=C double bond (unsaturated) | C≡C triple bond (unsaturated) |
| First member | Methane CH₄ | Ethene C₂H₄ | Ethyne C₂H₂ (acetylene) |
| Test | No reaction with Br₂ water | Decolourises bromine water | Decolourises bromine water |
| Main reaction type | Substitution (free radical) | Addition (electrophilic) | Addition |
Prefix: meth- (1C), eth- (2C), prop- (3C), but- (4C), pent- (5C), hex- (6C).
Suffix: -ane (single bond), -ene (double bond), -yne (triple bond).
Example: CH₃—CH₂—CH₃ = propane · CH₂=CH₂ = ethene · CH≡CH = ethyne
Bromine water test: Add bromine water (orange/yellow). Alkenes and alkynes decolourise it (addition reaction). Alkanes do NOT decolourise bromine water. This is the key test to distinguish saturated from unsaturated hydrocarbons — tested almost every year in WAEC and JAMB.
Functional Group Compounds
Alkanols contain the —OH group. Ethanol (C₂H₅OH) is the most commonly tested.
- Fermentation: C₆H₁₂O₆ → 2C₂H₅OH + 2CO₂ (glucose → ethanol + CO₂, in absence of oxygen, by yeast/zymase)
- Reactions of ethanol: Burns completely (combustion); oxidised to ethanoic acid (by acidified K₂Cr₂O₇ — orange to green); dehydrated by hot Al₂O₃ to ethene; reacts with Na metal to produce H₂; forms esters with carboxylic acids (esterification).
Alcohol + Carboxylic acid ⇌ Ester + Water
C₂H₅OH + CH₃COOH ⇌ CH₃COOC₂H₅ + H₂O (with conc. H₂SO₄ catalyst and heat)
Esters have sweet/fruity smells and are used in perfumes, flavourings and solvents.
Saponification is the hydrolysis of fats and oils (esters) by hot concentrated NaOH or KOH to produce soap (sodium/potassium salt of a long-chain fatty acid) and glycerol.
Fat/Oil + NaOH (conc., hot) → Soap + Glycerol
Soap vs Detergent: Soaps are sodium salts of fatty acids (from natural fats). Detergents are synthetic. Soap does NOT work well in hard water (forms scum). Detergents work in both hard and soft water.
How soap works: The soap molecule has a hydrophilic (water-loving) ionic head and a hydrophobic (water-hating/oil-loving) tail. The tail dissolves in grease; the head dissolves in water. This forms micelles that lift grease off surfaces and are rinsed away.
Chemistry & Industry
| Process | Product | Raw Materials | Conditions |
|---|---|---|---|
| Haber Process | Ammonia (NH₃) | N₂ (from air) + H₂ (from natural gas) | 450°C, 200 atm, iron catalyst |
| Contact Process | Sulphuric acid (H₂SO₄) | Sulphur + air + water | 450°C, V₂O₅ catalyst (for S→SO₂→SO₃→H₂SO₄) |
| Chlor-Alkali Process | Cl₂, NaOH, H₂ | Brine (NaCl solution) | Electrolysis |
| Solvay Process | Sodium carbonate (Na₂CO₃) | Salt, limestone, ammonia | Industrial temperatures |
| Petroleum refining | Petrol, kerosene, diesel, bitumen etc. | Crude oil | Fractional distillation |
| Cracking | Smaller alkanes + alkenes | Long-chain alkanes | High temperature ± catalyst; thermal or catalytic cracking |
Addition polymerisation: Monomers with C=C double bonds join together. No other products formed. Example: ethene → poly(ethene) [polythene]; chloroethene → PVC; propene → poly(propene).
Condensation polymerisation: Monomers join with loss of a small molecule (usually water). Examples: nylon (polyamide), terylene/polyester, proteins (amino acids), starch (glucose monomers).
Nigeria-specific chemistry: Nigeria produces crude oil (petroleum) — major export. Petrochemicals include plastics, fertilisers, lubricants. Fertilisers tested: NPK fertilisers contain nitrogen (N), phosphorus (P), potassium (K). Urea CO(NH₂)₂ is the most widely used nitrogen fertiliser.
You have now covered all the major WAEC and JAMB Chemistry topics. Take the timed CBT practice quiz to measure your score and get a personalised breakdown of your weak areas.