Cameroon GCE O-Level Chemistry cheat sheet

Matter is made of particles (atoms, ions, molecules) in constant random motion; the spaces between particles are largest in gases and smallest in solids.

Diffusion is the net movement of particles from high to low concentration caused by random motion; lighter gases diffuse faster, so rate is inversely proportional to the square root of molar mass.

Brownian motion: visible smoke or pollen particles jiggle because invisible fluid molecules collide with them unevenly, giving evidence for particle motion.

Solids: particles closely packed in fixed positions, vibrate only. Liquids: close but mobile, sliding past each other. Gases: far apart, moving fast and randomly.

Heating increases average kinetic energy of particles; cooling decreases it. Temperature is a measure of the average kinetic energy of particles, not the total energy.

Changes of state (melting, boiling, sublimation) break intermolecular forces, not covalent bonds; the particles themselves stay chemically unchanged.

A pure substance melts and boils at a fixed temperature; an impure substance melts over a range and at a lower temperature (melting point depression).

Compressibility decreases from gas to liquid to solid because the empty space between particles decreases in that order.

Particle diagrams: represent each species by a distinct symbol; a mixture shows different particles not chemically joined, a compound shows fixed ratios joined together.

Kinetic theory assumes gas particles are in random motion, have negligible volume relative to the container, and exert no forces on each other except during collisions.

Pressure of a gas arises from collisions of particles with container walls; increasing temperature or number of particles raises pressure if volume is fixed.

Boyle's law: at constant temperature, pV = constant, so p1V1 = p2V2.

Charles's law: at constant pressure, V/T = constant with T in kelvin, so V1/T1 = V2/T2.

Combined gas law: p1V1/T1 = p2V2/T2, with temperature always converted to kelvin (K = °C + 273).

Kelvin scale: 0 K = -273 °C; absolute zero is the temperature at which particle kinetic energy is theoretically minimum.

Boiling occurs when vapour pressure equals external pressure; reducing external pressure lowers the boiling point (as in vacuum distillation).

Evaporation occurs below the boiling point at the liquid surface; it is faster at higher temperature, larger surface area, and lower humidity.

Cooling curve: temperature falls until freezing begins, then stays constant while latent heat of fusion is released, then falls again.

Proton number (atomic number, Z) = number of protons; it determines the element. Nucleon number (mass number, A) = protons + neutrons.

In a neutral atom, electrons = protons. An ion's charge equals protons minus electrons.

Isotopes are atoms of the same element with the same proton number but different neutron numbers; they have identical chemical properties but different masses.

Relative atomic mass = sum of (isotope mass × fractional abundance); it is a weighted mean, not a simple average.

Electron shells fill 2, 8, 8, 18; the outer shell electron number equals the group number for main-group elements, and the number of shells equals the period number.

Group 1 (alkali metals) have one outer electron and form 1+ ions; Group 2 form 2+ ions; Group 7 (halogens) have seven outer electrons and form 1- ions.

Group 0 noble gases have full outer shells, are unreactive, and exist as single atoms.

Across a period, proton number increases, nuclear charge increases and atomic radius decreases; down a group, atomic radius increases and reactivity of metals increases.

Mass spectrometer output: peaks give m/z values and relative abundances; the most abundant isotope gives the tallest peak, not necessarily the heaviest.

Ionic bonding: electrostatic attraction between oppositely charged ions formed by transfer of electrons from a metal to a non-metal; giant lattice, high melting point, conducts when molten or aqueous.

Covalent bonding: a shared pair of electrons between non-metal atoms; simple molecular substances have weak intermolecular forces, low melting points, and do not conduct.

Giant covalent structures (diamond, silicon dioxide) have many strong covalent bonds, very high melting points, and do not conduct (except graphite).

Graphite conducts because each carbon uses only three of four outer electrons in bonding, leaving delocalised electrons between layers; layers slide, so it is soft.

Metallic bonding: positive ions in a lattice surrounded by a sea of delocalised electrons; explains conductivity, malleability and ductility.

Electron dot (Lewis) diagrams show outer-shell electrons only; each shared pair is one covalent bond, and a dative bond is a shared pair donated by one atom.

A molecule is a group of atoms held by covalent bonds; a lattice is a repeating three-dimensional arrangement of ions or atoms.

Solubility rule of thumb: ionic compounds dissolve in water because ion–dipole attractions overcome lattice energy; covalent molecules dissolve in non-polar solvents.

Alloys are harder than pure metals because different-sized atoms disrupt the layers, preventing them from sliding over each other.

Moles n = mass / molar mass; n = concentration (mol/dm³) × volume (dm³); n = gas volume (dm³) / 24 at room temperature (24 dm³ per mole).

Avogadro constant = 6.02 × 10²³ particles per mole; number of particles = moles × 6.02 × 10²³.

Percentage yield = (actual yield / theoretical yield) × 100; percentage purity = (mass of pure substance / mass of sample) × 100.

Empirical formula: divide each element's mass or percentage by its relative atomic mass, then divide by the smallest ratio. Molecular formula = empirical formula × n.

Balancing equations: atoms of each element and total charge must be equal on both sides; never change subscripts, only coefficients.

Limiting reactant: calculate moles of each reactant and divide by its coefficient; the smallest result identifies the limiting reactant.

Concentration in g/dm³ = concentration in mol/dm³ × molar mass.

Percentage of water of crystallisation = (mass of water / mass of hydrated compound) × 100.

Molar volume at s.t.p. (0 °C, 1 atm) is 22.4 dm³ per mole; at room conditions (25 °C) it is 24 dm³ per mole.

Acid: a proton (H⁺) donor; base: a proton acceptor; alkali: a soluble base that releases OH⁻ in water.

pH < 7 acidic, pH = 7 neutral, pH > 7 alkaline; pH = -log₁₀[H⁺] and [H⁺] = 10^(-pH).

Acid + metal → salt + hydrogen; acid + carbonate → salt + water + carbon dioxide; acid + base → salt + water (neutralisation).

Salt preparation: soluble salts by acid + excess insoluble base or carbonate, then filter, evaporate and crystallise; insoluble salts by precipitation (two soluble solutions mixed).

Titration calculation: moles of acid = concentration × volume; use the equation's mole ratio to find moles of base, then concentration = moles / volume.

Indicator choice: phenolphthalein (colourless to pink, 8.3–10.0) for strong acid–weak base or strong acid–strong base; methyl orange (red to yellow, 3.1–4.4) for strong acid–weak base.

Strong acids (HCl, HNO₃, H₂SO₄) ionise completely; weak acids (ethanoic, carbonic) ionise partially, giving lower [H⁺] at the same concentration.

Neutralisation is exothermic; the reaction of 1 mole of H⁺ with 1 mole of OH⁻ releases about 57 kJ.

Standard solution: a solution of accurately known concentration, prepared by dissolving a weighed mass in a volumetric flask and making up to the mark.

Exothermic: releases heat, ΔH negative, temperature rises; endothermic: absorbs heat, ΔH positive, temperature falls.

Bond breaking absorbs energy; bond forming releases energy; ΔH = energy of bonds broken − energy of bonds formed.

Rate increases with higher temperature (more particles exceed activation energy), higher concentration or pressure, larger surface area, and a catalyst (lowers activation energy).

Collision theory: reaction occurs only when particles collide with energy at least equal to the activation energy and in the correct orientation.

Le Chatelier's principle: a system at equilibrium shifts to oppose a change in concentration, pressure or temperature.

Increasing pressure shifts equilibrium towards the side with fewer gas molecules; increasing temperature shifts it in the endothermic direction.

A catalyst increases the rate of both forward and reverse reactions equally and does not change the position of equilibrium or the yield.

Equilibrium is dynamic: forward and reverse rates are equal and concentrations remain constant, but both reactions continue.

Oxidation is loss of electrons or gain of oxygen; reduction is gain of electrons or loss of oxygen; a redox reaction involves both.

Electrolysis: decomposition of an ionic compound by passing direct current through it molten or in aqueous solution.

Cathode (negative electrode): cations gain electrons (reduction). Anode (positive electrode): anions lose electrons (oxidation).

In aqueous solutions, the least reactive ion is discharged: at the cathode, metals below hydrogen in the reactivity series are deposited; otherwise hydrogen is produced.

At the anode with inert electrodes, halide ions are discharged before OH⁻; with sulfate or nitrate solutions, oxygen is produced from OH⁻.

Reactive electrodes (e.g. copper in copper sulfate) dissolve at the anode and plate onto the cathode, refining the metal.

Electrolysis of molten lead(II) bromide: lead at the cathode, bromine at the anode. Electrolysis of concentrated sodium chloride solution gives hydrogen, chlorine and sodium hydroxide.

Oxidation numbers: elements 0, simple ions equal to their charge, oxygen usually -2, hydrogen usually +1; the sum equals the overall charge.

Oxidising agent is reduced (gains electrons); reducing agent is oxidised (loses electrons).

Electroplating: the object to be plated is the cathode; the plating metal is the anode; the electrolyte contains the plating metal's ions.

Reactivity series (decreasing): potassium, sodium, calcium, magnesium, aluminium, zinc, iron, lead, copper, silver, gold.

Metals above carbon are extracted by electrolysis of molten compounds; metals below carbon are extracted by reduction with carbon.

Thermal decomposition of carbonates: metal carbonate → metal oxide + carbon dioxide; more reactive metals' carbonates are more stable to heat.

Metal + water: potassium and sodium give hydroxide + hydrogen; magnesium reacts with steam; iron and copper do not react with water.

Aluminium is extracted by electrolysis of alumina dissolved in molten cryolite; the cryolite lowers the melting point and saves energy.

Rusting requires both oxygen and water; salt accelerates rusting; prevention by painting, greasing, galvanising or sacrificial protection with zinc or magnesium.

Non-metals: oxygen supports combustion; chlorine bleaches damp litmus; hydrogen burns with a squeaky pop; nitrogen is relatively inert due to a strong triple bond.

Carbon dioxide turns limewater milky (calcium carbonate precipitate); sulfur dioxide decolourises acidified potassium manganate(VII).

Alloys: brass (copper + zinc), bronze (copper + tin), steel (iron + carbon), solder (lead + tin).

Alkanes CₙH₂ₙ₊₂ are saturated, undergo substitution with halogens in UV light, and burn: CH₄ + 2O₂ → CO₂ + 2H₂O.

Alkenes CₙH₂ₙ are unsaturated, contain a C=C double bond, and decolourise bromine water (orange to colourless) - the standard test.

Homologous series: same general formula, similar chemical properties, gradual change in physical properties, differ by CH₂.

Isomers have the same molecular formula but different structural formulae; they may have different boiling points.

Alcohols contain -OH; ethanol is made by fermentation of glucose with yeast or by hydration of ethene with steam and a catalyst.

Carboxylic acids (e.g. ethanoic acid) are weak acids that react with carbonates to give carbon dioxide and with alcohols to form esters.

Complete combustion gives CO₂ and H₂O; incomplete combustion gives carbon monoxide and soot; CO is toxic because it binds to haemoglobin.

Cracking breaks long-chain hydrocarbons into shorter alkenes and alkanes using heat and a catalyst; it increases petrol supply and produces alkenes for polymers.

Addition polymerisation joins alkene monomers with no other product; the polymer's repeat unit has the same atoms as the monomer.

Flame tests: lithium crimson, sodium yellow, potassium lilac, calcium brick-red, copper blue-green; clean the nichrome wire in concentrated HCl between tests.

Cations with NaOH: Cu²⁺ blue precipitate; Fe²⁺ green; Fe³⁺ red-brown; Zn²⁺ and Al³⁺ white precipitate soluble in excess NaOH; Ca²⁺ white precipitate insoluble in excess.

Cations with aqueous ammonia: Cu²⁺ blue precipitate soluble in excess to deep blue; Zn²⁺ white soluble in excess; Fe²⁺ and Fe³⁺ precipitates insoluble in excess.

Carbonate test: add dilute acid, bubbles of CO₂ turn limewater milky. Sulfate test: add dilute HCl then barium chloride → white precipitate.

Chloride test: add dilute nitric acid then silver nitrate → white precipitate; bromide gives cream, iodide gives yellow.

Ammonium test: warm with sodium hydroxide; ammonia gas turns damp red litmus blue.

Gas tests: hydrogen - lighted splint pops; oxygen - relights a glowing splint; chlorine - bleaches damp litmus; CO₂ - turns limewater milky.

Safety: wear goggles, tie back hair, never taste or smell chemicals directly, add acid to water (not water to acid) when diluting, and heat test tubes pointing away from people.

Bunsen burner: yellow flame is safe and luminous (air hole closed); blue roaring flame is hot and used for heating; never leave a lit burner unattended.

Fractional distillation of crude oil separates hydrocarbons by boiling point; fractions run from refinery gas (lowest) through petrol, naphtha, kerosene, diesel, fuel oil to bitumen (highest).

Longer-chain hydrocarbons have higher boiling points, higher viscosity and burn less cleanly.

Ethene test: shake with bromine water; orange-brown colour disappears. Alkanes do not decolourise bromine water in the dark.

Ethanol + ethanoic acid with a few drops of concentrated sulfuric acid catalyst gives ethyl ethanoate, an ester with a sweet smell.

Fermentation: C₆H₁₂O₆ → 2C₂H₅OH + 2CO₂, using yeast at about 30–35 °C; the ethanol concentration is limited to around 15% because yeast is killed.

Saturated means only single C–C bonds; unsaturated means at least one C=C double bond.

A fuel's efficiency depends on the energy released per gram and on complete combustion; hydrogen is a clean fuel because its only product is water.

General formula of alcohols is CₙH₂ₙ₊₁OH; of carboxylic acids is CₙH₂ₙ₊₁COOH.

Measuring volume: use a burette for titrations (read to the bottom of the meniscus, ±0.05 cm³), a pipette for a fixed aliquot, and a measuring cylinder for approximate volumes.

Weighing: use a balance accurate to 0.01 g or better; weigh by difference for hygroscopic or reactive solids.

Heating to constant mass: heat, cool in a desiccator, reweigh and repeat until two successive masses agree - used for determining water of crystallisation.

Filtration separates an insoluble solid from a liquid; crystallisation recovers a soluble solid; simple distillation separates a solvent from a solution.

Chromatography: the more soluble component travels further up the paper; Rf = distance moved by spot / distance moved by solvent front.

Test for water: anhydrous copper(II) sulfate turns blue; cobalt chloride paper turns from blue to pink.

Errors: systematic errors (zero error, wrongly calibrated apparatus) shift all readings; random errors scatter readings and are reduced by averaging.

When drawing apparatus, label all reagents and show the direction of gas flow; never draw a closed system for gas collection by displacement of air.

Hazard symbols: corrosive, toxic, flammable, oxidising, harmful/irritant, explosive, environmental hazard - each requires specific precautions.

Diluting concentrated sulfuric acid: add the acid slowly to water with stirring; the mixture becomes very hot and may spatter if reversed.

Bromine and chlorine are toxic; use them in a fume cupboard and never inhale directly.

Heating a liquid in a test tube: fill no more than one-third, hold with a holder, and point the mouth away from yourself and others.

Glassware: discard cracked or chipped apparatus; hot glass looks the same as cold glass, so allow it to cool before handling.

Spillage: neutralise acid spills with sodium carbonate and alkali spills with dilute acid, then mop up; report all spills to the teacher.

Waste disposal: never pour organic solvents or heavy-metal solutions down the sink; use the designated waste container.

Bunsen burner: light with the air hole closed, then adjust; the yellow flame is visible and safe, the blue flame is hotter but harder to see.