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Chemistry flashcards

210 cards covering every topic of the 9202 syllabus (Version 6.2, for exams May/June 2018 onwards). Study one topic or mix them all, turn each card, and mark yourself honestly.

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Topic 1: Atomic structure and the periodic table

Particles in a solid, a liquid and a gas: Solid: close together in a fixed, regular arrangement; vibrate. Liquid: close together, irregular; move past each other. Gas: far apart, random; move quickly in all directions.
What happens to the particles when a solid melts?: They gain enough energy to overcome some of the forces holding them in a fixed arrangement, and start to move past each other while staying close together.
What happens to the particles when a liquid boils or evaporates?: They gain enough energy to overcome the remaining forces between them and spread apart to fill the container.
Name the changes of state: solid → liquid, liquid → gas, gas → liquid, liquid → solid: Melting; boiling (or evaporating); condensing; freezing.
What do diffusion experiments (bromine in air, NH₃ and HCl, KMnO₄ in water) show?: That particles are in constant, random motion.
Ammonia and hydrogen chloride gases diffuse towards each other in a tube. What do you see?: A white ring of ammonium chloride where the two gases meet.
Element: A substance made of only one type of atom, represented by a chemical symbol.
Relative charge and relative mass of a proton, a neutron and an electron: Proton: +1, mass 1. Neutron: 0, mass 1. Electron: −1, very small mass.
Atomic number and mass number: Atomic number: the number of protons. Mass number: protons + neutrons.
Protons, electrons and neutrons in ²⁷₁₃Al: 13 protons, 13 electrons, 27 − 13 = 14 neutrons.
Isotopes: Atoms of the same element (same atomic number) with different numbers of neutrons, so different mass numbers.
Electron structure of sodium (atomic number 11): 2,8,1 — shells fill 2, then 8, then the remainder.
Relative atomic mass, Ar: An average mass of the atoms of an element, compared with carbon-12.
How did the model of the atom change from the plum pudding model?: Alpha-particle scattering showed the mass and positive charge are concentrated in a tiny nucleus, so the nuclear model replaced the plum pudding model. Electrons were then placed in shells, and the neutron was found later.
How are the elements arranged in the modern periodic table?: In order of atomic (proton) number, with elements with similar properties in the same columns (groups).
For Groups 1–7, what does the group number tell you?: The number of electrons in the outer shell.
How did Mendeleev’s table allow undiscovered elements to be predicted?: He grouped elements by similar properties and left gaps where no known element fitted, then predicted the properties of the missing elements from their groups.
Why are the noble gases (Group 0) unreactive?: They have a stable outer shell: 8 outer electrons (helium has 2), so they have no tendency to gain, lose or share electrons.
Why do Group 1 and Group 7 elements react?: To reach a stable, noble-gas-like outer shell: Group 1 atoms lose their one outer electron; Group 7 atoms gain one electron.

Topic 2: Structure, bonding and the properties of matter

Chemical bonding: Transferring or sharing electrons in the highest occupied energy levels (outer shells) so that atoms reach a noble gas electron arrangement.
Ion: An atom that has gained or lost electrons, so it carries an overall charge.
Ionic bonding: The strong electrostatic pull between oppositely charged ions, which works in every direction through a giant lattice.
Ion charges for Groups 1, 2, 3, 5, 6 and 7: Group 1: 1+. Group 2: 2+. Group 3: 3+. Group 5: 3−. Group 6: 2−. Group 7: 1−.
Charge and electron arrangement of the phosphide ion (P, atomic number 15): 3−; [2,8,8]³⁻ — the atom (2,8,5) gains three electrons.
Structure of sodium chloride: A regular cubic lattice: each Na⁺ is surrounded by six Cl⁻ and each Cl⁻ by six Na⁺, alternating in every direction.
Covalent bond: A pair of electrons shared by two atoms; it is a strong bond.
Shared pairs and non-bonding pairs in HCl, H₂O and NH₃: HCl: 1 shared, 3 non-bonding (on Cl). H₂O: 2 shared, 2 non-bonding (on O). NH₃: 3 shared, 1 non-bonding (on N).
Giant covalent structures (macromolecules) named in 9202: Diamond and silicon dioxide.
Structure and bonding in a metal: A giant structure of positive ions in a regular pattern, with delocalised outer electrons. Strong electrostatic attractions between the ions and the delocalised electrons hold it together.
Molecule vs giant structure: A molecule has a fixed, small number of atoms. A giant structure has no fixed size; its formula is only a ratio.
Explain why ionic compounds have high melting points: A giant ionic lattice has many strong electrostatic attractions between oppositely charged ions, acting in all directions. Overcoming them takes a large amount of energy.
When does an ionic compound conduct electricity, and why?: When molten or dissolved in water: the ions are free to move and carry charge. In the solid they are fixed in the lattice.
Melting and boiling points of simple molecular substances: Relatively low: they are gases, liquids or solids with relatively low melting and boiling points.
What is overcome when a simple molecular substance melts or boils?: Only the weak intermolecular forces between molecules. The covalent bonds are not broken.
Why do simple molecular substances not conduct electricity?: The molecules have no overall electric charge, so there are no charged particles to carry a current.
Why do giant covalent substances have very high melting points?: Every atom is held by strong covalent bonds, so a very large number of them must be broken. There are no separate molecules, so no intermolecular forces.
Why do metals conduct heat and electricity?: Their delocalised electrons can move through the whole structure, carrying charge and energy.
Substance Y melts at −210 °C and boils at −196 °C. What is its state at 25 °C?: Gas — 25 °C is above its boiling point.
Why can a carbon atom form four covalent bonds?: Its electron arrangement is 2,4, so it has four outer electrons to share.
Why is diamond hard?: Every carbon atom is locked in place by four strong covalent bonds pointing in all directions, so scratching it means breaking many strong bonds.
Why is graphite slippery?: There are no covalent bonds between the layers, only weak forces, so the layers slide.
Why does graphite conduct electricity?: One electron from each carbon atom is delocalised and moves through the structure.
Why does graphite still melt at a very high temperature?: Melting breaks the strong covalent bonds inside the layers.
Fullerenes: structure and the four uses named in 9202: Forms of carbon based on hexagonal rings, with different numbers of carbon atoms. Uses: drug delivery into the body, lubricants, catalysts, and nanotubes for reinforcing materials (e.g. tennis rackets).
Graphite vs metal conduction: Both use delocalised electrons. Graphite is not a metal; it is a giant covalent structure.
Nanoscale range: 1–100 nm, not "anything small". Of the order of a few hundred atoms.
Nanoparticle vs bulk: Same substance, different size, different properties, because of the far higher surface area to volume ratio.
Surface area to volume ratio of a cube with sides of 2 nm: Surface area = 6 × 2² = 24 nm². Volume = 2³ = 8 nm³. Ratio = 24 ÷ 8 = 3 nm⁻¹ (or 6 ÷ L).
Is a particle 6.5 × 10⁻⁸ m across a nanoparticle?: 6.5 × 10⁻⁸ ÷ 1 × 10⁻⁹ = 65 nm, which is within 1–100 nm, so yes.
Possible applications of nanoparticles named in 9202: New computers, new catalysts, new coatings, highly selective sensors, stronger and lighter construction materials, and new cosmetics such as suntan creams and deodorants.

Topic 3: Chemical changes

Why do metals conduct, and why can they be hammered into shape?: Conduct heat and electricity: delocalised electrons move through the structure. Hammered or bent: the layers of atoms slide over each other.
Alloy, and what steel is: Alloy: a mixture of at least two elements, at least one a metal (not a compound). Steel: iron mixed with carbon, sometimes with other metals too.
Why is copper used for wiring and plumbing?: Good conductor of heat and electricity; can be bent but is hard enough for pipes and tanks; does not react with water.
The 9202 reactivity series, most reactive first (with carbon): Potassium, sodium, lithium, calcium, magnesium, (carbon), zinc, iron, copper.
Copper wire is placed in silver nitrate solution. Ionic equation, and what is oxidised and reduced?: Cu + 2Ag⁺ → Cu²⁺ + 2Ag Copper atoms are oxidised (lose electrons); silver ions are reduced (gain electrons). Nitrate ions are spectators.
How is a metal’s extraction method chosen?: Very unreactive (gold): found as the metal. Below carbon (zinc, iron): heat the oxide with carbon (reduction). Above carbon (aluminium): electrolysis of the molten compound — high energy, so expensive.
Blast furnace: raw materials and the reduction step: Iron ore, coke, limestone and hot air. CO₂ + C → 2CO, then Fe₂O₃ + 3CO → 2Fe + 3CO₂.
Phytomining vs bioleaching: Phytomining: plants absorb copper compounds, then are burned; the ash contains copper compounds. Bioleaching: bacteria produce a leachate solution containing copper compounds.
Three reasons to recycle metals: Extraction uses limited resources, costs a lot of energy and damages the environment.
Thermal decomposition of copper carbonate: equation and colour change. Which carbonates do not decompose in a Bunsen flame?: CuCO₃ → CuO + CO₂; green → black. Sodium and potassium carbonates do not decompose at Bunsen temperatures.
6.20 g of CuCO₃ is heated to constant mass. Mass of copper oxide left? (Ar: Cu = 64, C = 12, O = 16): Mr CuCO₃ = 124, Mr CuO = 80. 6.20 × 80 ÷ 124 = 4.00 g
Electrolysis and electrolyte: Electrolysis: breaking down an ionic substance, molten or in solution, by passing an electric current through it. Electrolyte: the molten or dissolved ionic substance that is broken down.
Why must an ionic compound be melted or dissolved for electrolysis?: In the solid lattice the ions are fixed in place. Melting or dissolving frees them to move and carry charge.
Cathode and anode: charge, which ions go there, and what happens: Cathode: negative; positive ions move to it; reduction (gain of electrons). Anode: positive; negative ions move to it; oxidation (loss of electrons).
Aqueous solution: rules for the cathode and anode products: Cathode: hydrogen if the metal is more reactive than hydrogen; otherwise the metal (copper, silver). Anode: the halogen from a concentrated halide; otherwise (dilute halide, sulfate, nitrate) oxygen.
Required practical: copper sulfate solution with carbon electrodes. What is seen at each electrode?: Cathode: pink-brown copper coating, mass increases (Cu²⁺ + 2e⁻ → Cu). Anode: bubbles of oxygen, which relight a glowing splint (4OH⁻ → O₂ + 2H₂O + 4e⁻). The blue colour fades.
Electroplating: the cathode, anode and electrolyte, and three reasons for it: Cathode: the object. Anode: the plating metal. Electrolyte: a solution of ions of the plating metal. Reasons: appearance, durability, prevention of corrosion.
Aluminium manufacture: why cryolite, and why are the anodes replaced?: Aluminium oxide is dissolved in molten cryolite, which is molten at a much lower temperature, saving energy. The carbon anodes react with the oxygen formed (C + O₂ → CO₂), so they wear away.
Electrode half equations in aluminium manufacture: Negative electrode: Al³⁺ + 3e⁻ → Al Positive electrode: 2O²⁻ → O₂ + 4e⁻
Electrolysis of sodium chloride solution: products and uses: Cathode: hydrogen. Anode: chlorine. Solution: sodium hydroxide. Sodium hydroxide → soap; chlorine → bleach and plastics.

Topic 4: Chemical analysis

Pure substance (in chemistry): A single element or compound, with no other substance mixed in. Everyday "pure" (nothing added) is different: a "pure" natural product can still be a mixture.
How do impurities affect melting and boiling points?: A pure substance melts and boils at sharp, fixed values. Impurities lower the melting point and raise the boiling point, and make them happen over a range.
Why does purity matter for drugs and foodstuffs?: Drugs: the correct dose of the active compound, with no harmful impurities. Food: no harmful contaminants, and the contents match the label.
Choose the method: insoluble solid from a liquid; soluble solid from a solution; a liquid from a solution: Filtration; crystallisation; distillation.
Stationary and mobile phase in paper chromatography: Stationary phase: the paper. Mobile phase: the solvent.
Rf value, and a spot that moves 3.3 cm when the solvent front moves 5.5 cm: Rf = distance moved by substance ÷ distance moved by solvent, both from the baseline (to the centre of the spot). 3.3 ÷ 5.5 = 0.60. No units; never above 1.
How does paper chromatography separate the components of a mixture?: Each component is shared between the paper and the solvent. The more soluble it is in the solvent, the further it moves, so components with different solubilities separate.
Does one spot in one solvent prove a substance is pure?: No. Run it in a second solvent alongside a pure reference. One spot each time, level with the reference, is strong evidence.
Test for hydrogen: A lighted splint near the gas: a pop (2H₂ + O₂ → 2H₂O).
Test for oxygen: A glowing splint in the gas relights.
Test for carbon dioxide, and why it works: Bubble the gas through limewater (calcium hydroxide solution): it turns cloudy white. Ca(OH)₂ + CO₂ → CaCO₃ + H₂O — insoluble calcium carbonate forms.
Two tests for ammonia: Damp red litmus paper turns blue. Hydrogen chloride from concentrated hydrochloric acid held near it gives white smoke of ammonium chloride (NH₃ + HCl → NH₄Cl).
Test for chlorine: Damp blue litmus paper turns red, then is bleached white.
Flame test colours for Li⁺, Na⁺, K⁺, Ca²⁺ and Ba²⁺: Li⁺: crimson. Na⁺: yellow. K⁺: lilac. Ca²⁺: red. Ba²⁺: green.
Flame test practical: how is the wire cleaned, and what can spoil the result?: Dip a nichrome loop in dilute hydrochloric acid and heat, repeating until the flame shows no colour. Sodium contamination gives a strong yellow that can hide other colours (especially lilac); a mixture of ions may show one colour only.
Sodium hydroxide solution: precipitates with Cu²⁺, Fe²⁺ and Fe³⁺: Cu²⁺: blue. Fe²⁺: green (slowly turns brown in air). Fe³⁺: brown.
Al³⁺, Ca²⁺ and Mg²⁺ all give a white precipitate with NaOH. How do you tell them apart?: Add excess NaOH: only aluminium hydroxide dissolves. For Ca²⁺ vs Mg²⁺ use a flame test: calcium gives red; magnesium gives no characteristic colour.
Test for carbonate ions: Add dilute acid: fizzing. Pass the gas into limewater: it turns cloudy white (carbon dioxide).
Test for halide ions, and the precipitate colours: Acidify with dilute nitric acid, then add silver nitrate solution. Cl⁻: white. Br⁻: cream. I⁻: yellow.
Test for sulfate ions, and why sulfuric acid is not used: Add dilute hydrochloric acid, then barium chloride solution: a white precipitate of barium sulfate. Sulfuric acid contains sulfate ions, so it would always give a precipitate.
A solution gives a brown precipitate with NaOH and a white precipitate with acidified BaCl₂. Name and formula?: Iron(III) sulfate, Fe₂(SO₄)₃ — two Fe³⁺ balance three SO₄²⁻.

Topic 5: Acids, bases and salts

Base and alkali: Base: a metal oxide or metal hydroxide. Alkali: a soluble hydroxide (a base that dissolves in water). Copper(II) oxide is a base, not an alkali.
Salt: The compound formed when the hydrogen of an acid is replaced by a metal (or ammonium) ion.
Salt endings from hydrochloric, nitric and sulfuric acid: Hydrochloric → chloride. Nitric → nitrate. Sulfuric → sulfate.
acid + metal → acid + metal oxide or hydroxide → acid + ammonia →: salt + hydrogen salt + water an ammonium salt (no water)
Why is a solution of ammonia alkaline?: Ammonia reacts with water to form hydroxide ions: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻.
pH ranges and universal indicator colours: 0–2 very acidic (red); 3–6 weakly acidic (orange/yellow); 7 neutral (green); 8–11 weakly alkaline (blue); 12–14 very alkaline (purple).
Ionic equation for neutralisation: H⁺(aq) + OH⁻(aq) → H₂O(l). H⁺ comes from the acid, OH⁻ from the alkali; the other ions are spectators.
Balance: Al(OH)₃ + HNO₃ → Al(NO₃)₃ + H₂O: Al(OH)₃ + 3HNO₃ → Al(NO₃)₃ + 3H₂O
Making a soluble salt from a metal or an insoluble base: Warm the acid; stir in the solid until some is left over (excess); filter off the excess; crystallise the filtrate.
Which metals are unsuitable for the metal + acid route, and why?: Potassium, sodium, lithium: too reactive — violent and dangerous. Copper, silver, gold: not reactive enough — no reaction with dilute hydrochloric or sulfuric acid.
Making a salt from an alkali and an acid: Pipette the alkali into a flask with indicator; add acid from a burette until the colour just changes; repeat with the same volumes and no indicator; crystallise.
Crystallisation: method: Heat to evaporate some water until saturated (crystals form on a cooled glass rod); cool slowly; filter, rinse with a little cold water and dry between filter papers. Do not heat to dryness.
Making an insoluble salt such as barium sulfate: Mix two solutions, e.g. barium nitrate and potassium sulfate; filter; wash the residue with distilled water; dry. Ba(NO₃)₂(aq) + K₂SO₄(aq) → BaSO₄(s) + 2KNO₃(aq)
How is precipitation used in water treatment?: To remove unwanted ions: calcium ions from hard water as calcium carbonate; metal ions from effluent as insoluble hydroxides; phosphate as an insoluble phosphate.

Topic 6: Quantitative chemistry

Why is mass conserved in a chemical reaction?: Atoms are rearranged, not created or destroyed, so the total mass of the products equals the total mass of the reactants.
A metal carbonate is heated in an open crucible and its mass falls. Explain.: Carbon dioxide gas is produced and escapes. In a closed system the total mass would be unchanged.
A metal is heated in air and its mass increases. Explain.: A gas from the air (oxygen) is a reactant, so its mass is added to the product.
What does 2Mg + O₂ → 2MgO tell you quantitatively?: 2 moles of magnesium react with 1 mole of oxygen to form 2 moles of magnesium oxide. The coefficients, not the subscripts, give the reacting ratio.
Relative formula mass, Mr: The sum of the relative atomic masses of all the atoms in the formula.
Why are relative atomic masses not whole numbers?: They are weighted averages of the masses of an element’s isotopes, taking their abundances into account.
Equation linking moles, mass and Mr: moles = mass (g) ÷ Mr, so mass = moles × Mr.
Three steps of a reacting-mass calculation: 1. Convert the known mass to moles. 2. Use the equation’s mole ratio to find moles of the wanted substance. 3. Convert back to mass. Never use the mass ratio directly.
Mass of MgO from burning 4.8 g of magnesium (2Mg + O₂ → 2MgO; Ar: Mg = 24, O = 16): 4.8 ÷ 24 = 0.2 mol Mg → 0.2 mol MgO. 0.2 × 40 = 8.0 g
Finding an empirical formula from masses or percentages: Divide each by its Ar, divide all results by the smallest, then scale to whole numbers. Molecular formula = empirical formula × (Mr ÷ empirical formula mass).
A compound is 40.0% C, 6.7% H and 53.3% O, with Mr 180. Empirical and molecular formulae?: C 3.33, H 6.7, O 3.33 → 1 : 2 : 1, CH₂O (mass 30). 180 ÷ 30 = 6, so C₆H₁₂O₆.
Why might you not obtain the calculated mass of product?: The reaction may not go to completion (it is reversible); some product may be lost when it is separated; some reactants may react in unexpected ways. No atoms are lost.
The mole and the Avogadro constant: One mole contains 6.02 × 10²³ particles (the Avogadro constant). Its mass in grams equals the relative formula mass.
Number of particles from moles: number of particles = moles × 6.02 × 10²³
Volume of one mole of any gas at room temperature and pressure: 24 dm³ (24 000 cm³). volume (dm³) = moles × 24.
Why do gas volume ratios equal the mole ratios in an equation?: Equal volumes of any gases at the same temperature and pressure contain equal numbers of molecules.
Volume of CO₂ at rtp from 2.50 g of CaCO₃ with excess HCl (CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂; Mr CaCO₃ = 100): 2.50 ÷ 100 = 0.0250 mol CaCO₃ → 0.0250 mol CO₂. 0.0250 × 24 = 0.600 dm³
Concentration in g/dm³ and in mol/dm³: g/dm³ = mass ÷ volume. mol/dm³ = moles ÷ volume. g/dm³ = mol/dm³ × Mr. Volume must be in dm³ (cm³ ÷ 1000).
Moles of HCl in 25.0 cm³ of 2.00 mol/dm³ hydrochloric acid: 2.00 × (25.0 ÷ 1000) = 0.0500 mol
25.0 cm³ of 0.0500 mol/dm³ H₂SO₄ neutralises 20.0 cm³ of NaOH (H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O). Concentration of NaOH?: H₂SO₄: 1.25 × 10⁻³ mol → NaOH: 2.50 × 10⁻³ mol (1 : 2). 2.50 × 10⁻³ ÷ 0.0200 = 0.125 mol/dm³
Which titres are used to calculate the mean?: Only concordant titres (within 0.10 cm³ of each other). Leave out the rough titre.

Topic 7: Trends within the periodic table

Group 1 metal + water: products, and the pattern equation: The metal hydroxide (an alkaline solution) and hydrogen. 2M + 2H₂O → 2MOH + H₂
Lithium, sodium and potassium with water: observations: Lithium: floats, steady fizzing. Sodium: floats, melts into a ball, moves about quickly. Potassium: floats; the hydrogen ignites with a lilac flame.
Compounds of Group 1 metals with non-metals: Ionic compounds containing the M⁺ ion; white solids that dissolve to give colourless solutions.
Halogens: form of the element, ion charge and trends down the group: Diatomic molecules (F₂, Cl₂, Br₂, I₂). They form 1− halide ions with metals. Down the group: reactivity decreases; melting and boiling points increase.
Appearance of fluorine, chlorine, bromine and iodine at room temperature: Fluorine: pale yellow gas. Chlorine: pale yellow-green gas. Bromine: red-brown liquid. Iodine: grey-black solid.
Chlorine water + lithium bromide solution: equation and what you see: Cl₂ + 2LiBr → 2LiCl + Br₂ (ionic: Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂). The solution turns yellow-orange.
Explain why Group 1 metals get more reactive down the group: More energy levels: the outer electron is further from the nucleus with more shielding, so the attraction is weaker and the electron is lost more easily.
Explain why halogens get less reactive down the group: More energy levels: the outer level is further from the nucleus with more shielding, so the attraction for an incoming electron is weaker and an electron is gained less easily.
Where are the transition metals, and what three properties do many of them have?: In the centre of the periodic table, between Groups 2 and 3. Many form ions with different charges, form coloured compounds and are useful as catalysts.
Transition metals compared with Group 1 metals: Higher melting points (except mercury) and higher densities; stronger and harder; much less reactive, so they react far less vigorously with water and oxygen.
Which transition metal is the melting-point exception?: Mercury: liquid at room temperature, melting at −39 °C. The exception is for melting point only — its density is still high.
Charge on the vanadium ion in V₂O₅, and the compound’s name: Five O²⁻ give 10−, shared by two V ions: V⁵⁺. Vanadium(V) oxide.
Name CuBr and give the formula of iron(III) nitrate: CuBr: copper(I) bromide (Br⁻ is 1−, so Cu⁺). Iron(III) nitrate: Fe(NO₃)₃.
Is every transition metal compound coloured?: No — "many", not all. Zinc compounds and anhydrous copper(II) sulfate are white.
Transition metal catalysts in the Haber and Contact processes: Haber process: iron. Contact process (stage 2, SO₂ → SO₃): vanadium(V) oxide, V₂O₅.

Topic 8: The rate and extent of chemical change

Ways to calculate a mean rate of reaction, with units: Mass of reactant used ÷ time (g/s); volume of gas formed ÷ time (cm³/s); moles ÷ time (mol/s). For a time to a visible change, 1 ÷ time (s⁻¹).
A gas syringe reads 21 cm³ at 20 s and 57 cm³ at 50 s. Mean rate over this interval?: (57 − 21) ÷ (50 − 20) = 36 ÷ 30 = 1.2 cm³/s
Collision theory and activation energy: Particles react only when they collide with sufficient energy. Activation energy is the minimum energy particles need for a reaction to happen.
Why does raising the temperature increase the rate?: Particles move faster, so they collide more frequently and more energetically. A greater proportion of collisions have at least the activation energy.
How do concentration, pressure and surface area increase the rate?: Concentration and pressure: more particles in the same volume. Surface area: more particles exposed. Each gives more frequent collisions — but not more energetic ones.
Catalyst: what it does and why industry uses one: Alters the rate of a reaction and is not used up (same mass at the end); it does not change the amount of product. Industry: faster production, lower temperatures and pressures (less energy), and it can be reused, reducing costs.
A product curve levels off. What does that mean, and what does a faster run with the same reactants look like?: The reaction has stopped: a reactant is used up. A faster run starts steeper, levels off sooner, at the same height.
Equilibrium (9202 definition): In a closed system, the forward and backward reactions occur at exactly the same rate. Both reactions continue; the amounts stay constant but are usually not equal.
Effect of raising the temperature on an equilibrium: The yield of the endothermic direction increases; the yield of the exothermic direction decreases. (Lowering the temperature does the opposite.)
Effect of increasing the pressure on an equilibrium of gases: It favours the reaction producing the fewest molecules of gas. If both sides have the same number of molecules, the yield does not change.
PCl₅(g) ⇌ PCl₃(g) + Cl₂(g). Effect of increasing the pressure?: 1 molecule on the left, 2 on the right, so the backward reaction is favoured: more PCl₅ remains.
Does a catalyst change the yield at equilibrium?: No. It speeds up the forward and backward reactions equally, so equilibrium is reached sooner but the amounts are the same.
Four things industry balances when choosing conditions: Yield, rate, cost and safety. For an exothermic forward reaction a low temperature gives the best yield but is too slow, so a compromise temperature is used.
Haber process: equation, raw materials and conditions: N₂(g) + 3H₂(g) ⇌ 2NH₃(g). Nitrogen from the air; hydrogen from natural gas or other sources. Iron catalyst, about 450 °C, about 200 atmospheres.
Haber process: how is ammonia separated, and what happens to unreacted gases?: Cooling: ammonia condenses at a much higher temperature than nitrogen or hydrogen, so it liquefies and is removed. Unreacted nitrogen and hydrogen are recycled.
Why is about 450 °C used in the Haber process?: The forward reaction is exothermic, so a lower temperature would give a higher yield but would be too slow. 450 °C is a compromise.
Why not use a pressure much higher than 200 atmospheres in the Haber process?: Higher pressure gives a higher yield (4 → 2 molecules) and faster rate, but needs stronger, more expensive equipment and more energy, and is more dangerous.
Why are fertilizers made from ammonia needed?: Crops need nitrogen to make proteins, and most cannot take it from the air; harvesting removes nutrients; the population is rising and farmland is limited, so fertilizers maximise food yields.
The three stages of the contact process: 1. S + O₂ → SO₂ 2. 2SO₂ + O₂ ⇌ 2SO₃ (exothermic) 3. H₂O + SO₃ → H₂SO₄
Contact process stage 2: conditions, and why atmospheric pressure is used: Vanadium(V) oxide catalyst, around 450 °C, atmospheric pressure. The yield is already high, so the extra cost of high pressure is not justified.
Oxidation and reduction in terms of oxygen: Oxidation: gain of oxygen by a substance. Reduction: loss of oxygen from a substance.
Oxidation and reduction in terms of electrons: Oxidation is loss of electrons; reduction is gain of electrons (OIL RIG).
In Fe₂O₃ + 3CO → 2Fe + 3CO₂, what is oxidised and what is reduced?: Fe₂O₃ loses oxygen: reduced. CO gains oxygen: oxidised.
Zinc in lead(II) nitrate solution: ionic equation and half equations: Zn + Pb²⁺ → Zn²⁺ + Pb Zn → Zn²⁺ + 2e⁻ (oxidation) Pb²⁺ + 2e⁻ → Pb (reduction)
Is MgO + 2HCl → MgCl₂ + H₂O a redox reaction?: No. The ions only change partners; no particle gains or loses electrons and nothing gains oxygen. Neutralisation, precipitation and carbonate decomposition are not redox.
Fe²⁺ → Fe³⁺ + e⁻: oxidation or reduction?: Oxidation — an electron is lost, so the ion becomes more positive.

Topic 9: Energy changes

Exothermic and endothermic reactions: Exothermic: transfers energy to the surroundings, which get hotter; ΔH negative. Endothermic: takes in energy from the surroundings, which get colder; ΔH positive.
Three types of exothermic reaction, and two everyday uses: Combustion, many oxidation reactions (e.g. rusting) and neutralisation. Uses: self-heating cans and hand warmers.
A type of endothermic reaction, and an everyday use: Thermal decomposition (e.g. calcium carbonate → calcium oxide + carbon dioxide). Use: sports injury packs that go cold (e.g. ammonium nitrate dissolving in water).
Hydrated copper sulfate ⇌ anhydrous copper sulfate + water: colours and energy changes: Heating: blue → white, endothermic. Adding water: white → blue and it gets hot, exothermic. The same amount of energy is transferred each way.
Q = mcΔT: what each symbol means, with units: Q: energy transferred (J). m: mass of water or solution heated (g) — not the fuel. c: specific heat capacity (4.2 J/g °C for water). ΔT: temperature change (°C).
0.50 g of a fuel heats 100 g of water by 15.0 °C. Energy released per gram? (c = 4.2 J/g °C): Q = 100 × 4.2 × 15.0 = 6300 J = 6.30 kJ. 6.30 ÷ 0.50 = 12.6 kJ/g
Why are calorimetry results usually lower than the true value?: Energy heats the container, thermometer and air; fuels burn incompletely; fuel evaporates from the wick; slow reactions lose more heat before the highest temperature is reached.
Reaction profile: how are activation energy and ΔH read, and what does a catalyst change?: Activation energy = peak − reactant level. ΔH = product level − reactant level. A catalyst gives a lower peak with the same start, finish and ΔH.
Bond breaking and bond making: energy and the overall sign: Breaking bonds takes energy in; making bonds gives energy out. Energy change = bonds broken − bonds formed: negative (exothermic) when more is released making bonds.
CO + H₂O → CO₂ + H₂. Bond energies: C≡O 1077, O–H 464, C=O 805, H–H 436 kJ/mol. Energy change?: Broken: 1077 + 2 × 464 = 2005. Formed: 2 × 805 + 436 = 2046. 2005 − 2046 = −41 kJ/mol (exothermic).
Chemical cell, and why it stops producing a potential difference: It contains reactants that react to transfer energy electrically. It produces a pd until a reactant is used up — not because it "runs out of electricity".
Fuel cell, and how it differs from a chemical cell: It produces electricity from the reaction of a fuel with oxygen, both supplied from outside, so it works as long as they keep being supplied. It is refuelled, not replaced.
Hydrogen–oxygen fuel cell: overall reaction, waste product and uses: 2H₂ + O₂ → 2H₂O, without burning. Water is the only waste product. Uses: cars and spacecraft.
Why do fuel cells cause less pollution only "when in use"?: Only water is released at the vehicle, but making the hydrogen may use fossil fuels or electricity from them, so pollution can be released elsewhere.
Two disadvantages of storing hydrogen as a fuel: It is a very low-density gas, so a useful mass has a huge volume and must be kept at high pressure in strong, heavy tanks (or cooled to a liquid); it is highly flammable and leaks can explode.
Why might a hydrogen engine release a pollutant that a fuel cell does not?: The engine burns hydrogen in air at high temperature, so oxides of nitrogen can form. The fuel cell does not burn it.
Mass of water made from 3.0 kg of hydrogen in a fuel cell (H = 1, O = 16): 3000 ÷ 2 = 1500 mol H₂ → 1500 mol H₂O. 1500 × 18 = 27 000 g = 27 kg (always 9 × the mass of hydrogen).

Topic 10: Organic chemistry

Crude oil vs the hydrocarbons in it: Crude oil is a mixture; each hydrocarbon in it is a compound.
How does fractional distillation separate crude oil?: The oil is vaporised and fed into a column that is hotter at the base than the top. Each vapour rises until it reaches a level cooler than its boiling point, then condenses and is tapped off. It is continuous.
Saturated vs unsaturated: Saturated: only single C–C bonds. Unsaturated: at least one C=C.
General formulae of alkanes and alkenes, and the formula of the alkane with 20 carbons: Alkanes: CₙH₂ₙ₊₂. Alkenes: CₙH₂ₙ. C₂₀H₄₂
How do boiling point, viscosity and flammability change as hydrocarbon molecules get bigger?: Boiling point rises, viscosity increases and flammability decreases — so short-chain hydrocarbons make better fuels.
Complete and incomplete combustion of ethane: Complete: 2C₂H₆ + 7O₂ → 4CO₂ + 6H₂O. Incomplete (oxygen short): CO or soot forms, e.g. 2C₂H₆ + 5O₂ → 4CO + 6H₂O. Water still forms.
Acid rain vs global dimming vs climate change: the pollutants: Acid rain: SO₂ (sulfur in the fuel) and NOₓ (air reacting at high temperature). Global dimming: particulates. Climate change: CO₂.
Making ethanol: hydration of ethene vs fermentation: Hydration: C₂H₄ + H₂O → C₂H₅OH; phosphoric acid, about 300 °C, high pressure; fast, continuous, pure; non-renewable raw material. Fermentation: yeast enzymes, 20–35 °C; slow, batch, dilute; renewable sugar.
Biofuels: renewability, land use and carbon footprint: Renewable: crops regrow. Land use: fuel crops can take land from food or lead to forest clearing. Carbon footprint: growing plants absorb CO₂, but farming, processing and transport add emissions, so it is reduced, not removed.
Cracking: what it is and the two methods: Thermal decomposition of large hydrocarbons into smaller alkanes and alkenes. Vaporise them, then pass over a hot catalyst, or mix with steam and heat to a very high temperature.
Fractional distillation vs cracking: Fractional distillation: physical separation by boiling point. Cracking: a chemical reaction that breaks molecules.
Required practical: testing for a C=C double bond: Shake with bromine water. Alkene: orange → colourless. Alkane: stays orange.
Monomer and polymer: Monomer: a small molecule that can join with many others. Polymer: a very large molecule formed from many monomers joined together.
Addition polymerisation of an alkene: what happens?: One bond of each C=C opens and the freed bonds link neighbouring monomers into a long chain. Nothing else is made. Ethene → poly(ethene); propene → poly(propene).
Give the monomer of the chain unit –(CHF–CF₂)n–: CHF=CF₂ — remove the bonds through the brackets and the n, and put a double bond between the two carbons.
LD vs HD poly(ethene): Same monomer, different conditions and catalysts. LD: highly branched chains, loose packing, weaker forces — flexible, lower density. HD: mainly unbranched, close packing, stronger forces — rigid, higher density.
Thermosoftening vs thermosetting polymers: Thermosoftening: separate chains held by weak intermolecular forces; soften and melt on heating. Thermosetting: chains joined by strong covalent cross-links; do not melt.
Why are non-biodegradable polymers a problem, and what does biodegradable mean?: Biodegradable: can be broken down by microbes. Many polymers are not, so they persist in landfill and as litter that harms wildlife.
Polymer + cornstarch bag vs biodegradable plastic made from cornstarch: Polymer + cornstarch: microbes digest the starch so the bag falls apart, but polymer fragments may remain. Cornstarch plastic: the whole material can be broken down by microbes.
Functional groups of alcohols, carboxylic acids and esters: Alcohols: –OH (names end -ol). Carboxylic acids: –COOH (end -oic acid). Esters: –COO– (two words, the second ending -oate).
Alcohols: effect on water, reaction with sodium, and combustion: Dissolve to give a neutral solution (pH 7). With sodium: fizzing, hydrogen given off. Burn completely to carbon dioxide and water.
Complete combustion of ethanol: C₂H₅OH + 3O₂ → 2CO₂ + 3H₂O
Two ways to oxidise ethanol to ethanoic acid: A chemical oxidising agent (e.g. acidified potassium dichromate), or microbial action using oxygen from the air — why opened wine turns sour. Ethanoic acid is the main acid in vinegar.
Ethanoic acid + potassium carbonate: balanced equation: 2CH₃COOH + K₂CO₃ → 2CH₃COOK + H₂O + CO₂ (fizzing; the gas turns limewater milky).
Why does 0.1 mol/dm³ ethanoic acid have a higher pH than 0.1 mol/dm³ nitric acid?: Ethanoic acid is a weak acid: it only partly ionises (CH₃COOH ⇌ CH₃COO⁻ + H⁺), so it gives fewer H⁺ ions. Nitric acid ionises completely.
Making ethyl ethanoate, and two properties and uses of esters: Ethanol + ethanoic acid ⇌ ethyl ethanoate + water, with an acid catalyst. Esters are volatile with distinctive smells, used in flavourings and perfumes.

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Written for OxfordAQA International GCSE Chemistry (9202), syllabus 9202 (Version 6.2, for exams May/June 2018 onwards).

Cards by Marlbridge Academic Team; 23 restate definitions from the 9202 revision notes. This deck was reviewed by Nouman Ahmed on 2026-11-05. Report an error.

All 9202 notes, practice and topics