Free flashcards · Cambridge O Level · 5070

Chemistry flashcards

273 cards covering every topic of the 5070 syllabus (2026-2028). Study one topic or mix them all, turn each card, and mark yourself honestly.

Study mode needs JavaScript. Every card is listed below, so you can still cover the answers and test yourself.

All 273 cards

For a quick read-through or printing. Cover the right-hand side to test yourself.

Topic 1: States of matter

Particles in a solid: arrangement and movement: Close together in a regular arrangement. They vibrate about fixed positions.
Particles in a liquid: arrangement and movement: Close together in an irregular arrangement. They move around, sliding past each other.
Particles in a gas: arrangement and movement: Far apart in a random arrangement. They move quickly in all directions.
Melting point: The temperature at which a solid becomes a liquid.
Boiling point: The temperature at which a liquid becomes a gas throughout, forming bubbles.
Evaporation: Liquid to gas at the surface only, at any temperature below the boiling point.
Name the change of state: gas → liquid, and liquid → solid: Gas → liquid: condensation. Liquid → solid: freezing (solidification).
A gas at constant pressure is heated. What happens to its volume? A gas at constant temperature is compressed (pressure increased). What happens to its volume?: Heating: the volume increases. Increasing the pressure: the volume decreases.
Why is a heating curve flat while a substance melts or boils?: The energy supplied goes into overcoming the forces of attraction between particles, not into raising their kinetic energy, so the temperature stays constant.
Explain why a gas expands when heated at constant pressure: The particles gain kinetic energy and move faster, so they hit the container walls harder and more often. The gas expands until the pressure is back to its original value.
Diffusion: 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.
Effect of temperature on the rate of diffusion: Higher temperature → faster diffusion, because the particles move faster.
Effect of relative molecular mass (Mr) on the rate of diffusion of a gas: Lower Mr → faster diffusion. Lighter particles move faster at the same temperature.
Ammonia and hydrogen chloride diffuse from opposite ends of a tube. Where does the white ring form, and why?: Nearer the hydrogen chloride end. NH₃ (Mr 17) is lighter than HCl (Mr 36.5), so it diffuses faster and travels further. The white solid is ammonium chloride.

Topic 2: Atoms, elements and compounds

Element: A substance made of only one type of atom.
Compound: A substance made of two or more elements chemically combined.
Mixture: Two or more substances that are not chemically combined.
How can you tell a mixture from a compound?: A mixture can be separated by physical methods (such as filtration or distillation), and each part keeps its own properties. A compound has different properties from its elements and can only be split by a chemical reaction.
Relative charge and relative mass of a proton, a neutron and an electron: Proton: charge +1, mass 1. Neutron: charge 0, mass 1. Electron: charge −1, mass 1/1836 (negligible).
Proton number (atomic number): The number of protons in the nucleus of an atom.
Mass number (nucleon number): The total number of protons and neutrons in the nucleus of an atom.
Electronic configuration of sodium (proton number 11): 2,8,1
Electronic configuration of chlorine (proton number 17): 2,8,7
What do the group number and the period number tell you about an atom’s electrons?: Group number (I–VII) = number of outer-shell electrons. Period number = number of occupied electron shells.
Isotopes: Atoms of the same element with the same number of protons but different numbers of neutrons.
How many neutrons are in chlorine-35 and chlorine-37? (proton number 17): Chlorine-35: 18 neutrons. Chlorine-37: 20 neutrons.
Why do isotopes of an element have the same chemical properties?: They have the same number of electrons, and so the same number of outer-shell electrons (the same electronic configuration).
Chlorine is 75% chlorine-35 and 25% chlorine-37. Calculate its relative atomic mass.: (75 × 35 + 25 × 37) ÷ 100 = 35.5
Ion: An atom or group of atoms that has gained or lost electrons, so it carries a charge.
Cation and anion: Cation: a positive ion (electrons lost). Anion: a negative ion (electrons gained).
How an ionic bond forms: The metal loses electrons (forming a cation) and the non-metal gains them (forming an anion). Both reach a full outer shell.
Ionic bond: A strong electrostatic attraction between oppositely charged ions.
How do sodium and chlorine atoms form sodium chloride?: Each sodium atom loses one electron to form Na⁺. Each chlorine atom gains one electron to form Cl⁻. The oppositely charged ions attract.
Electrical conductivity of an ionic compound: Poor when solid. Good when molten or dissolved in water.
Structure of an ionic compound such as sodium chloride: A giant lattice: a regular arrangement of alternating positive and negative ions.
Explain why ionic compounds have high melting points: The strong electrostatic attractions between oppositely charged ions act throughout the giant lattice, so a lot of energy is needed to overcome them.
Explain why an ionic compound conducts when molten but not when solid: In the solid the ions are held in fixed positions in the lattice. When molten (or in solution) the ions are free to move and carry charge.
How a covalent bond forms: Two non-metal atoms share a pair of electrons, so each reaches a full outer shell.
How many covalent bonds are there in a molecule of methane, CH₄, and of water, H₂O?: Methane: four single C–H bonds. Water: two single O–H bonds.
Properties of simple molecular substances: Low melting and boiling points. Poor electrical conductivity.
Explain why simple molecular substances have low melting and boiling points: Only the weak intermolecular forces between molecules have to be overcome, which needs little energy. The covalent bonds inside the molecules are not broken.
Explain why simple molecular substances do not conduct electricity: They have no ions or free electrons, so there are no mobile charged particles.
Bonding in carbon dioxide and in nitrogen: CO₂: two C=O double bonds (O=C=O). N₂: one N≡N triple bond.
Structure of diamond: A giant covalent lattice in which each carbon atom is bonded to four others in a tetrahedral arrangement.
Structure of graphite: Layers of hexagonal rings. Each carbon atom is bonded to three others, with weak forces between the layers and one delocalised electron per atom.
Why is diamond used in cutting tools?: It is very hard, because every atom is held by strong covalent bonds in a rigid three-dimensional lattice.
Why is graphite used as a lubricant and as an electrode?: Lubricant: the layers slide over each other easily. Electrode: its delocalised electrons move through the structure, so it conducts electricity.
Structure of silicon(IV) oxide, SiO₂: A giant covalent structure similar to diamond: each silicon atom is bonded to four oxygen atoms and each oxygen atom to two silicon atoms.
Metallic bonding: An electrostatic attraction between a lattice of positive metal ions and a “sea” of delocalised electrons.
How metallic bonding forms: Metal atoms lose their outer electrons into a delocalised sea, leaving positive ions.
Explain why metals conduct electricity: The delocalised electrons can move through the structure and carry charge.
Explain why metals are malleable and ductile: The layers of positive ions can slide over each other without the metallic bonding being broken.

Topic 3: Stoichiometry

Molecular formula: The number and type of each atom in one molecule.
State symbols: (s) solid, (l) liquid, (g) gas, (aq) aqueous — dissolved in water.
Formulae of the sulfate, nitrate, carbonate, hydroxide and ammonium ions: SO₄²⁻, NO₃⁻, CO₃²⁻, OH⁻, NH₄⁺
Empirical formula: The simplest whole-number ratio of the different atoms in a compound.
Use the ion charges to write the formula of aluminium oxide (Al³⁺, O²⁻): Al₂O₃. The lowest common multiple of 3 and 2 is 6, so 2 Al³⁺ balance 3 O²⁻.
Formula of calcium nitrate (Ca²⁺, NO₃⁻): Ca(NO₃)₂ — brackets because two nitrate ions are needed.
Ionic equation for the neutralisation of an acid by an alkali: H⁺(aq) + OH⁻(aq) → H₂O(l)
Relative atomic mass, Ar: The average mass of the isotopes of an element compared with 1/12 of the mass of an atom of carbon-12.
Relative molecular mass, Mr: The sum of the relative atomic masses of all the atoms in a molecule. (For an ionic compound: relative formula mass.)
Mr of carbon dioxide, CO₂ (Ar: C = 12, O = 16): 12 + 2 × 16 = 44
Mr of calcium carbonate, CaCO₃ (Ar: Ca = 40, C = 12, O = 16): 40 + 12 + 3 × 16 = 100
Two units for the concentration of a solution: g/dm³ and mol/dm³
The mole: The amount of substance that contains 6.02 × 10²³ particles (the Avogadro constant).
Equation linking moles, mass and molar mass: moles = mass (g) ÷ molar mass (g/mol)
Molar gas volume at room temperature and pressure (rtp): 24 dm³ (24 000 cm³) per mole of any gas.
Equation linking moles, concentration and volume: moles = concentration (mol/dm³) × volume (dm³). Convert cm³ to dm³ by dividing by 1000.
Percentage yield: actual yield ÷ theoretical yield × 100. It can never be more than 100%.
Mass of 0.25 mol of CaCO₃ (Mr = 100): 0.25 × 100 = 25 g

Topic 4: Electrochemistry

Electrolysis: The decomposition of an ionic compound, when molten or in aqueous solution, by the passage of an electric current.
Anode and cathode: Anode: the positive electrode. Cathode: the negative electrode. (PANIC — Positive is Anode, Negative Is Cathode.)
Products of electrolysing molten lead(II) bromide with inert electrodes: Cathode: lead. Anode: bromine (red-brown vapour).
Products of electrolysing concentrated aqueous sodium chloride with inert electrodes: Cathode: hydrogen. Anode: chlorine.
Products of electrolysing dilute sulfuric acid with inert electrodes: Cathode: hydrogen. Anode: oxygen. Volume ratio hydrogen : oxygen = 2 : 1.
Electroplating: the cathode, the anode and the electrolyte: Cathode: the object to be plated. Anode: the plating metal. Electrolyte: a solution containing ions of the plating metal.
What happens at each electrode, in terms of electrons?: Cathode: reduction — ions gain electrons. Anode: oxidation — ions lose electrons. (AN OIL RIG CAT)
Half-equations for electrolysing molten lead(II) bromide: Cathode: Pb²⁺ + 2e⁻ → Pb Anode: 2Br⁻ → Br₂ + 2e⁻
Aqueous solution: which is discharged at the cathode, the metal or hydrogen?: The less reactive of the two. A metal below hydrogen in the reactivity series (Cu, Ag) is deposited; otherwise hydrogen is given off.
Products of electrolysing dilute aqueous sodium chloride with inert electrodes: Cathode: hydrogen. Anode: oxygen (the halide is not concentrated enough to be discharged).
Copper(II) sulfate solution electrolysed with copper electrodes: Copper is deposited at the cathode. The copper anode dissolves (Cu → Cu²⁺ + 2e⁻) instead of oxygen being released.
Hydrogen–oxygen fuel cell: Uses hydrogen and oxygen to produce electricity. The only product is water.
Advantages of a hydrogen fuel cell over a petrol engine: Only water is produced, so no CO₂ at the point of use. More efficient than combustion.
Disadvantages of a hydrogen fuel cell: Hydrogen is hard to store and transport, highly flammable and expensive, and producing it may itself release CO₂.

Topic 5: Chemical energetics

Exothermic reaction: Transfers thermal energy to the surroundings, so the temperature of the surroundings rises.
Endothermic reaction: Takes in thermal energy from the surroundings, so the temperature of the surroundings falls.
On a reaction pathway diagram, how can you tell a reaction is exothermic?: The products are at a lower energy level than the reactants.
Sign of ΔH for exothermic and endothermic reactions: Exothermic: ΔH is negative. Endothermic: ΔH is positive.
Activation energy, Eₐ: The minimum energy that colliding particles must have to react.
Bond breaking and bond making: which is endothermic?: Bond breaking is endothermic; bond making is exothermic. ΔH = energy to break bonds − energy released making bonds.

Topic 6: Chemical reactions

Physical change: No new substance is formed, and the change is usually easy to reverse (for example, melting ice).
Chemical change: A new substance is formed, and the change is usually hard to reverse (for example, burning).
Signs that a chemical change has happened: A new substance forms — often shown by a colour change, a gas given off (effervescence), a precipitate, or a temperature change.
Four ways to increase the rate of a reaction: Raise the temperature; increase the concentration (or the pressure of a gas); increase the surface area (use smaller pieces); add a catalyst.
Catalyst: A substance that increases the rate of a reaction and is unchanged in mass and chemically at the end of the reaction.
Two ways to follow the rate of a reaction that gives off a gas: Measure the volume of gas with a gas syringe at regular times, or measure the loss in mass on a balance.
Collision theory: why does raising the temperature increase the rate?: Particles move faster, so collisions are more frequent, and a greater proportion of collisions have energy equal to or greater than the activation energy.
Collision theory: why does increasing concentration increase the rate?: There are more particles per unit volume, so collisions are more frequent.
Collision theory: why does increasing surface area increase the rate?: More particles are exposed, so collisions are more frequent.
How does a catalyst increase the rate?: It provides an alternative pathway with a lower activation energy.
What does the symbol ⇌ mean?: The reaction is reversible: it can go in both directions.
Heating hydrated copper(II) sulfate, then adding water: Heating: blue hydrated crystals → white anhydrous powder. Adding water: white → blue again.
Colours of hydrated and anhydrous cobalt(II) chloride: Hydrated: pink. Anhydrous: blue.
Features of a dynamic equilibrium: The forward and reverse reactions happen at the same rate, so the concentrations stay constant. It needs a closed system.
Effect of increasing pressure on an equilibrium involving gases: The position of equilibrium shifts towards the side with fewer gas molecules.
Haber process: equation and conditions: N₂ + 3H₂ ⇌ 2NH₃ 450 °C, 200 atm, iron catalyst.
Haber process: sources of nitrogen and hydrogen: Nitrogen: the air. Hydrogen: methane.
Contact process: equation and conditions: 2SO₂ + O₂ ⇌ 2SO₃ 450 °C, 200 kPa (2 atm), vanadium(V) oxide catalyst.
Why is 450 °C used in the Haber process rather than a lower temperature?: It is a compromise: a lower temperature would give a higher yield (the forward reaction is exothermic) but the rate would be too slow to be economic.
Why is 200 atm used in the Haber process rather than a much higher pressure?: Higher pressure would push the equilibrium further towards ammonia, but very high pressures are expensive to generate and maintain and are a safety risk. 200 atm is a compromise.
Contact process: sources of sulfur dioxide and oxygen: Sulfur dioxide: burning sulfur or roasting sulfide ores. Oxygen: the air.
Oxidation and reduction (oxygen definition): Oxidation: gain of oxygen. Reduction: loss of oxygen.
Redox reaction: A reaction in which oxidation and reduction happen at the same time.
What does the Roman numeral in iron(III) oxide tell you?: The oxidation number of iron in the compound is +3.
Oxidation and reduction (electron definition): Oxidation is loss of electrons; reduction is gain of electrons. (OIL RIG)
Oxidising agent: Causes oxidation in something else, so is itself reduced and gains electrons.
Reducing agent: Causes reduction in something else, so is itself oxidised and loses electrons.
Oxidation number: what does an increase and a decrease mean?: Increase in oxidation number: oxidation. Decrease: reduction.

Topic 7: Acids, bases and salts

pH of acidic, neutral and alkaline solutions: Acidic: below 7. Neutral: 7. Alkaline: above 7.
Colours of litmus, methyl orange and thymolphthalein in acid and in alkali: Litmus: red in acid, blue in alkali. Methyl orange: red in acid, yellow in alkali. Thymolphthalein: colourless in acid, blue in alkali.
Which ions do aqueous acids and aqueous alkalis contain?: Acids: H⁺ ions. Alkalis: OH⁻ ions.
Alkali: A soluble base. All alkalis are bases, but not all bases are alkalis.
acid + metal → acid + base → acid + carbonate →: salt + hydrogen salt + water salt + water + carbon dioxide
What is made when an alkali is warmed with an ammonium salt?: Ammonia gas (plus a salt and water).
Acid (proton definition): A proton (H⁺) donor.
Base (proton definition): A proton (H⁺) acceptor.
Strong acid: Fully dissociated in water. Examples: HCl, HNO₃, H₂SO₄.
Weak acid: Only partially dissociated in water. Examples: ethanoic, citric and carbonic acids.
Acidic oxides: Oxides of non-metals, such as SO₂ and CO₂. They react with bases and alkalis.
Basic oxides: Oxides of metals, such as CuO and CaO. They react with acids.
Amphoteric oxides: Oxides that react with both acids and bases to form a salt and water. Examples: aluminium oxide, Al₂O₃, and zinc oxide, ZnO.
Making a soluble salt from an acid and an insoluble base (or carbonate or metal): Add the solid in excess to the acid, filter off the leftover solid, then crystallise the filtrate.
When is titration used to make a salt?: When both reactants are soluble (acid + alkali), so there is no excess solid to filter off. Repeat without the indicator to get a pure salt.
Which common salts are always soluble?: All sodium, potassium and ammonium salts, and all nitrates.
Exceptions to the solubility rules for chlorides and sulfates: Chlorides: soluble except silver and lead. Sulfates: soluble except barium, calcium and lead.
Solubility of carbonates and hydroxides: Mostly insoluble, except sodium, potassium and ammonium (calcium hydroxide is partially soluble).
Hydrated and anhydrous: Hydrated: a substance chemically combined with water (e.g. CuSO₄·5H₂O). Anhydrous: containing no water.
Making an insoluble salt: Precipitation: mix two solutions containing its ions, then filter, wash the precipitate with distilled water and dry it.
Water of crystallisation: The water molecules chemically combined within the crystal structure of a hydrated salt.

Topic 8: The Periodic Table

How is the Periodic Table arranged?: In order of increasing proton number. Columns are groups; rows are periods.
How does character change across a period?: From metallic on the left to non-metallic on the right.
Why do elements in the same group have similar chemical properties?: They have the same number of outer-shell electrons.
Physical properties of lithium, sodium and potassium: Relatively soft metals with low density and low melting points.
Group I: trends down the group: Reactivity increases. Melting point decreases. Density generally increases.
Group I metal + water →: metal hydroxide + hydrogen e.g. 2Na + 2H₂O → 2NaOH + H₂
Colour and state of chlorine, bromine and iodine at room temperature: Chlorine: pale yellow-green gas. Bromine: red-brown liquid. Iodine: grey-black solid.
Group VII: trends down the group: Reactivity decreases. Colour gets darker. Density and melting point increase.
Chlorine water is added to potassium bromide solution. What happens?: Chlorine is more reactive, so it displaces bromine: chlorine + potassium bromide → potassium chloride + bromine. The solution turns orange.
Halogen molecules: Diatomic: Cl₂, Br₂, I₂.
Properties of transition elements: High densities and high melting points. They form coloured compounds, and they and their compounds often act as catalysts.
Transition elements have ions with variable oxidation numbers. Give an example.: Iron forms Fe²⁺ (iron(II)) and Fe³⁺ (iron(III)) compounds. 5070 states the general property; iron is the usual example.
Transition elements compared with Group I metals: density, melting point and reactivity: Transition elements: high density, high melting point, low reactivity. Group I metals: low density, low melting point, very high reactivity.
Compounds of Group I metals and of transition elements: colour: Group I compounds are white or colourless. Transition element compounds are coloured.
Ion charges: Group I metals compared with transition elements: Group I metals always form 1+ ions. Transition elements form ions with variable charges, e.g. Fe²⁺ and Fe³⁺.
Noble gases (Group VIII): physical form: Unreactive, monatomic gases.
Why are the noble gases unreactive?: They have a full outer shell of electrons, so they do not need to lose, gain or share electrons.
Electronic configurations of neon and argon, and what they explain: Neon is 2,8 and argon is 2,8,8. Each has a full outer shell, so its atoms have no tendency to lose, gain or share electrons: the gases are unreactive.
Monatomic: Existing as single, separate atoms rather than as molecules. The noble gases are monatomic, unlike the diatomic halogens.
A use of helium, of neon and of argon: Helium: balloons (low density, non-flammable). Neon: lighting. Argon: lamps and welding (an inert atmosphere).

Topic 9: Metals

General physical properties of metals: Good conductors of heat and electricity, malleable, ductile, and usually high melting and boiling points.
General physical properties of non-metals: Poor conductors of heat and electricity, brittle when solid, and usually low melting and boiling points.
metal + dilute acid → metal + oxygen →: salt + hydrogen metal oxide
Aluminium: two uses and the property behind each: Aircraft parts — low density. Food containers — resistance to corrosion.
Copper: a use and the properties behind it: Electrical wiring — good electrical conductor and ductile.
Aluminium in overhead electrical cables: which properties?: Low density and good electrical conductivity.
Copper in electrical wiring: which properties, and what does ductile mean?: Good electrical conductivity and ductility. Ductile means it can be drawn out into wires.
Alloy: A mixture of a metal with other elements.
What is brass made from?: Copper and zinc.
Stainless steel: composition and a use: Iron mixed with chromium, nickel and carbon. Used for cutlery, because it is hard and resists corrosion.
How do the hardness and strength of an alloy compare with the pure metal?: Alloys can be harder and stronger than the pure metals.
Explain why an alloy is harder than the pure metal: The different-sized atoms disrupt the regular layers, so the layers cannot slide over each other as easily.
The reactivity series (most to least reactive), with carbon and hydrogen: Potassium, sodium, calcium, magnesium, aluminium, (carbon), zinc, iron, (hydrogen), copper, silver, gold.
Which metals react with cold water, and which only with steam?: Cold water: potassium, sodium, calcium. Steam: magnesium (very slow with cold water), zinc and iron.
Will magnesium displace zinc from zinc sulfate solution?: Yes. Magnesium is more reactive than zinc, so it gives up electrons more readily and displaces it.
Why does aluminium seem unreactive?: It is covered by a thin, protective layer of aluminium oxide.
What two things are needed for iron to rust?: Water and oxygen.
Rust: Hydrated iron(III) oxide.
Barrier methods of rust prevention: Painting, greasing (oiling) and coating with plastic — they keep out water and oxygen.
Sacrificial protection: Attaching a more reactive metal (such as zinc or magnesium) to iron. It corrodes instead of the iron.
Galvanising, and why it beats tin plating: Coating iron with zinc. It is a barrier and also sacrificial, so iron is still protected if scratched. Tin is less reactive than iron, so scratched tin plate rusts faster.
How a metal’s position in the reactivity series decides its extraction method: More reactive than carbon: electrolysis. Less reactive than carbon: reduction by heating with carbon.
Raw materials of the blast furnace: Hematite (iron ore), coke (carbon), limestone and hot air.
What reduces iron(III) oxide in the blast furnace?: Carbon monoxide.
Main ore of aluminium, and how aluminium is extracted: Bauxite. Aluminium is extracted by electrolysis.
Equation for the reduction of iron(III) oxide in the blast furnace: Fe₂O₃ + 3CO → 2Fe + 3CO₂
How limestone removes impurities in the blast furnace: CaCO₃ → CaO + CO₂, then CaO + SiO₂ → CaSiO₃ (slag).
Role of cryolite in aluminium extraction: Aluminium oxide is dissolved in molten cryolite, which lowers the operating temperature and improves conductivity.

Topic 10: Chemistry of the environment

Two chemical tests for water: Anhydrous copper(II) sulfate turns from white to blue. Anhydrous cobalt(II) chloride turns from blue to pink.
Physical test for pure water: It melts at exactly 0 °C and boils at exactly 100 °C.
Stages of water treatment, in order: 1. Sedimentation and filtration — remove solids. 2. Carbon — removes tastes and odours. 3. Chlorination — kills bacteria.
NPK fertilisers: the three elements: Nitrogen, phosphorus and potassium.
Role of nitrogen, phosphorus and potassium in plants: Nitrogen: protein synthesis, leaf growth. Phosphorus: root growth. Potassium: flowering, fruit, disease resistance.
Two types of compound used as nitrogen fertilisers: Ammonium salts and nitrates.
Composition of clean, dry air: About 78% nitrogen and 21% oxygen. The rest is mainly argon, with a little carbon dioxide.
Carbon monoxide: source and effect: Incomplete combustion of carbon-containing fuels. Toxic — it reduces the blood’s ability to carry oxygen.
Sulfur dioxide: source and effect: Burning fossil fuels that contain sulfur impurities. Causes acid rain.
Two greenhouse gases and their sources: Carbon dioxide: complete combustion of fuels, respiration. Methane: livestock, decomposition of vegetation, rice paddies.
What does a catalytic converter in a car remove?: Oxides of nitrogen (and carbon monoxide), cutting acid rain.
Catalytic converter equation removing CO and NO together: 2CO + 2NO → 2CO₂ + N₂
Symbol equation for photosynthesis: 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂

Topic 11: Organic chemistry

Functional group: An atom or group of atoms that determines the chemical properties of a homologous series.
Homologous series: A family of compounds with the same functional group and similar chemical properties.
Saturated: Contains only single carbon–carbon bonds (e.g. alkanes).
Unsaturated: Contains at least one carbon–carbon double or triple bond, most commonly C=C (e.g. alkenes).
Homologous series: the full description: Same functional group, same general formula, each member differs by CH₂, a gradual trend in physical properties, similar chemical properties.
Structural isomers: Compounds with the same molecular formula but different structural formulae.
Name endings for alkanes, alkenes, alcohols and carboxylic acids: Alkane: -ane. Alkene: -ene. Alcohol: -ol. Carboxylic acid: -oic acid.
Name C₂H₅OH and CH₃COOH: Ethanol and ethanoic acid.
Name the two unbranched alkenes with formula C₄H₈: But-1-ene and but-2-ene.
Name the ester formed from propanoic acid and methanol: Methyl propanoate — the alcohol part is named first.
Three fossil fuels, and the main constituent of natural gas: Coal, natural gas and petroleum. Natural gas is mainly methane.
How is petroleum separated?: By fractional distillation, into fractions with different boiling point ranges.
Going UP the fractionating column: chain length, boiling point, volatility, viscosity: Chain length decreases, boiling point is lower, volatility is higher, viscosity is lower.
Uses of: refinery gas, gasoline, naphtha, kerosene: Refinery gas: heating and cooking. Gasoline/petrol: car fuel. Naphtha: chemical feedstock. Kerosene/paraffin: jet fuel.
Uses of: diesel oil, fuel oil, lubricating oil, bitumen: Diesel oil/gas oil: diesel engines. Fuel oil: ships and home heating. Lubricating oil: lubricants, waxes, polishes. Bitumen: making roads.
General formula of the alkanes: CₙH₂ₙ₊₂
Reactivity of alkanes: Generally unreactive, apart from combustion. They are saturated hydrocarbons.
Products of the complete combustion of an alkane: Carbon dioxide and water.
Reaction of methane with chlorine: type, condition and equation: Substitution, in ultraviolet light. CH₄ + Cl₂ → CH₃Cl + HCl
General formula of the alkenes: CₙH₂ₙ
Cracking: Breaking larger alkane molecules into smaller alkanes, alkenes and hydrogen, using a high temperature and a catalyst.
Why is cracking done?: There is more demand for short-chain fractions such as petrol than is supplied, and the alkenes made are useful (for example, to make polymers).
Test to tell an alkene from an alkane: Add bromine water. Alkene: orange → colourless. Alkane: stays orange.
Addition reactions of ethene: catalyst and product with hydrogen and with steam: Hydrogen: nickel catalyst → ethane. Steam: acid catalyst → ethanol.
Two ways of making ethanol: Fermentation of glucose with yeast. Catalytic addition of steam to ethene.
Conditions for fermentation: 25–35 °C, yeast, absence of oxygen.
Conditions for making ethanol from ethene and steam: 300 °C, 6000 kPa (60 atm), acid catalyst.
Two uses of ethanol: As a solvent and as a fuel.
Fermentation compared with hydration of ethene: Fermentation: renewable raw material, but slow, batch process, dilute impure product. From ethene: fast, continuous, pure product, but uses a non-renewable raw material and lots of energy.
Complete combustion of an alcohol: products and two examples: Carbon dioxide and water. C₂H₅OH + 3O₂ → 2CO₂ + 3H₂O 2C₃H₇OH + 9O₂ → 6CO₂ + 8H₂O (propan-1-ol)
Ethanoic acid + a metal: Salt + hydrogen. 2CH₃COOH + Mg → (CH₃COO)₂Mg + H₂ (magnesium ethanoate)
Ethanoic acid + a base: Salt + water. CH₃COOH + NaOH → CH₃COONa + H₂O (sodium ethanoate)
Ethanoic acid + a carbonate: Salt + water + carbon dioxide. 2CH₃COOH + Na₂CO₃ → 2CH₃COONa + H₂O + CO₂
Two ways of making ethanoic acid: Oxidising ethanol with acidified potassium manganate(VII), or by bacterial oxidation (fermentation of ethanol in air).
Making an ester: Carboxylic acid + alcohol, with an acid catalyst → ester + water. Ethanoic acid + ethanol → ethyl ethanoate.
Why does ethanoic acid react more slowly than hydrochloric acid of the same concentration?: Ethanoic acid is a weak acid: it is only partially dissociated in water, so there are fewer H⁺ ions.
Do other carboxylic acids react like ethanoic acid? Give an example.: Yes — methanoic and propanoic acids give methanoates and propanoates. 2C₂H₅COOH + Mg → (C₂H₅COO)₂Mg + H₂ (magnesium propanoate)
Polymer: A large molecule built up from many small units called monomers.
Problems with disposing of plastics: Most are non-biodegradable, so they fill landfill and harm aquatic life; burning them can release toxic gases.
Addition polymerisation: Many monomers with a C=C double bond join together. Nothing else is formed. Ethene → poly(ethene).
Polyamide (nylon): A dicarboxylic acid + a diamine → amide link –CONH–, with water lost.
Polyester (PET): A dicarboxylic acid + a diol → ester link –COO–, with water lost.
Proteins as polymers: Natural polyamides formed from amino acid monomers.

Topic 12: Experimental techniques and chemical analysis

Apparatus to measure time, temperature, mass and volume: Time: stopwatch. Temperature: thermometer. Mass: balance. Volume: burette, volumetric pipette, measuring cylinder, gas syringe.
Solvent: A substance that dissolves a solute.
Solute: A substance that is dissolved in a solvent.
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: A substance that remains after evaporation, distillation, filtration or any similar process.
Filtrate: A liquid or solution that has passed through a filter.
Apparatus for an acid–base titration: Burette, volumetric pipette, conical flask and a suitable indicator.
When do you stop adding from the burette, and how many runs do you need?: Stop when the indicator just changes colour (the end-point). Repeat until you have concordant titres.
Why not use universal indicator in a titration?: It changes colour gradually across the pH range, so it gives no sharp end-point. Use methyl orange or thymolphthalein.
Why is the baseline on a chromatogram drawn in pencil, above the solvent?: Pencil does not dissolve and run. Above the solvent so the spots do not dissolve into the solvent in the beaker.
How does a chromatogram show that a substance is pure?: A pure substance gives a single spot.
Rf value: distance moved by the substance ÷ distance moved by the solvent front, both measured from the baseline.
A spot moves 4.5 cm and the solvent front 7.5 cm. What is the Rf?: 4.5 ÷ 7.5 = 0.6
How are colourless substances seen on a chromatogram?: By spraying the paper with a locating agent.
Filtration separates…: An insoluble solid from a liquid. The solid is the residue; the liquid is the filtrate.
Crystallisation: what it separates and the method: A dissolved solid from its solution. Heat to evaporate some solvent until saturated, then cool; filter off the crystals, wash with a little cold solvent and dry.
Simple distillation separates…: A solvent from a dissolved solid (for example, water from salt water). The solvent evaporates and is condensed in a condenser.
Fractional distillation separates…: Miscible liquids with different boiling points (for example, ethanol and water). Uses a fractionating column; the liquid with the lower boiling point distils first.
Separating one soluble solid from an insoluble one: Add a suitable solvent to dissolve the soluble one, then filter.
How do melting and boiling points show whether a substance is pure?: A pure substance has a sharp, fixed melting point and boiling point. Impurities lower the melting point and raise the boiling point, and make them happen over a range.
Tests for hydrogen and oxygen: Hydrogen: a lighted splint gives a squeaky pop. Oxygen: relights a glowing splint.
Tests for carbon dioxide and ammonia: Carbon dioxide: turns limewater milky. Ammonia: turns damp red litmus paper blue.
Tests for chlorine and sulfur dioxide: Chlorine: bleaches damp litmus paper. Sulfur dioxide: turns acidified potassium manganate(VII) from purple to colourless.
Aqueous sodium hydroxide: precipitate colours for Cu²⁺, Fe²⁺ and Fe³⁺: Cu²⁺: light blue. Fe²⁺: green. Fe³⁺: red-brown. All insoluble in excess.
Zn²⁺ and Al³⁺ both give a white precipitate that dissolves in excess NaOH. How do you tell them apart?: Use aqueous ammonia: the zinc precipitate dissolves in excess; the aluminium precipitate does not.
Test for halide ions, and the precipitate colours: Acidify with dilute nitric acid, add aqueous silver nitrate. Cl⁻: white. Br⁻: cream. I⁻: yellow.
Test for sulfate ions: Acidify with dilute nitric acid, add aqueous barium nitrate: white precipitate.
Test for carbonate ions: Add dilute acid: effervescence, and the gas turns limewater milky.
Flame test colours: Li⁺, Na⁺, K⁺, Ca²⁺, Ba²⁺, Cu²⁺: Li⁺: red. Na⁺: yellow. K⁺: lilac. Ca²⁺: orange-red. Ba²⁺: light green. Cu²⁺: blue-green.

Flashcards by topic

Written for Cambridge O Level Chemistry (5070), syllabus 5070 (2026-2028).

Cards by Marlbridge Academic Team; 40 restate definitions from the 5070 revision notes. This deck was reviewed by Farhat ul Ain Sehgal on 2026-11-05. Report an error.

All 5070 notes, practice and topics