AQA · GCSE

化学

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Handouts · ⁨讲义⁩ · IGCSE Chemistry · ⁨IGCSE 化学⁩ (12)

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Lessons · ⁨课程⁩ →Vocabulary practice · ⁨词汇练习⁩ →

4.1 · Atomic structure and the periodic table

A substance containing only one type of atom defined by proton number.
Symbols use an uppercase first letter and, where present, a lowercase second letter: Co is cobalt, whereas CO is carbon monoxide. For the first twenty elements learn H hydrogen, He helium, Li lithium, Be beryllium, B boron, C carbon, N nitrogen, O oxygen, F fluorine, Ne neon, Na sodium, Mg magnesium, Al aluminium, Si silicon, P phosphorus, S sulfur, Cl chlorine, Ar argon, K potassium and Ca calcium. Group 1 also includes Rb rubidium, Cs caesium and Fr francium; Group 7 includes F fluorine, Cl chlorine, Br bromine, I iodine and At astatine. Use the supplied periodic table for other specified elements, including Fe iron, Cu copper and Zn zinc. Very radioactive francium and astatine are names to interpret from the table, not school reaction specimens.
Read a formula by identifying each element symbol before applying its subscript. A subscript applies to the symbol immediately before it, or to an entire bracketed group. Name MgO magnesium oxide, NaCl sodium chloride and Ca(OH)₂ calcium hydroxide. Compare diagrams of separate atoms, molecules of one element and compound particles; keep particle identity distinct from the number of particles.
A number before a chemical formula multiplying the entire formula.
Begin with the actual substance formulae and make an element-count table. Change coefficients to equalize the count for each element. Never change MgO to MgO₂ just to balance oxygen: that would identify a different substance. For hydrogen burning, 2H₂ + O₂ → 2H₂O conserves four hydrogen atoms and two oxygen atoms. State symbols, when required, distinguish solid (s), liquid (l), gas (g) and aqueous solution (aq).
Use coloured counters to represent atoms and construct the reactants from the stated formulae. Rearrange the same counters into product particles, then record the smallest whole-number coefficients. Confirm each element separately. Models count atoms and bonds schematically; they do not show the detailed collision mechanism or actual particle size.
An equation showing electron loss or gain for one oxidation or reduction process.
In the magnesium half equation, the right-hand total charge is +2 − 2 = 0, matching neutral Mg. In copper reduction, +2 − 2 = 0 on the left, matching neutral copper. Neutralization can be represented as H⁺(aq) + OH⁻(aq) → H₂O(l); the net charge is zero on each side. Electron numbers must match when combining half equations so no free electrons remain in the overall reaction.
First write the relevant ions and products, then balance atoms and charge separately. For a full aqueous equation, expand soluble strong electrolytes into their ions and cancel only identical species present on both sides. Retain solid precipitates, water and other unchanged molecular forms as appropriate. Use supplied ion charges; the equation is not a licence to invent a charge.
The insoluble solid retained by a filter.
For sand mixed with salt, add water and stir to dissolve the salt, filter the insoluble sand, then gently evaporate some water from the filtrate. Stop concentration before boiling dry and allow the solution to cool; filter and dry the crystals. Washing the sand removes adhering solution, and washing crystals with a little cold suitable solvent removes some surface impurity while limiting dissolution.
Use actual teacher-supervised school separation equipment, eye protection and a risk-assessed heat source. Fold and support the paper in a funnel, collect filtrate in a clean vessel and transfer solution to an evaporating basin. Use a safe end-point such as crystals forming in a cooled test drop. Hot glass can appear cold. Record mass only after the recovered material is dry.
Liquid collected after vapour is condensed during distillation.
Paper chromatography separates soluble components according to their different distributions between the moving solvent and the paper. Draw the baseline in pencil above the solvent level, place a small sample spot and let solvent rise. Components that spend more time in the mobile phase generally travel farther under these conditions. Compare spots with known standards using the same solvent and paper; a single spot alone does not prove purity under every possible method.
For school-supervised distillation, use a risk-assessed apparatus with an open receiving path, never a sealed heated system. Cooling water enters the lower condenser connection and leaves the upper one. Avoid flames with flammable solvents. In chromatography, keep the lid on the vessel as appropriate, mark the solvent front before it dries and measure all distances from the same baseline. Interpretation of Rf is developed in Chemical analysis, 4.8.1.3.
An atomic model with mass and positive charge concentrated in a tiny central nucleus.
The nuclear model replaced diffuse positive charge with a dense nucleus and electrons outside it. Bohr added electrons at specific distances or energy levels; his theoretical predictions agreed with observations. Later work identified positively charged protons. Chadwick provided evidence for uncharged neutrons in the nucleus about twenty years after the nuclear idea became accepted. A scientific model can develop further without every earlier observation becoming wrong.
Use supplied scattering observations or a computer-free diagram, not a classroom radioactive source. Match each observation to an inference and compare it with the diffuse-charge prediction. Distinguish observed trajectories from inferred structure. A marble analogue illustrates deflection but cannot establish electric charge or reproduce quantum behaviour. Detailed experimental work behind Bohr and Chadwick is outside this specification requirement.
The number of protons in an atom’s nucleus.
Net relative charge = proton number − electron number. Losing electrons produces a positive ion; gaining electrons produces a negative ion. Ordinary chemical ion formation changes electrons, not the nucleus. An atom with eleven protons is sodium even when it has ten electrons; its net charge is +1. A model with ten protons is neon, regardless of a similar electron arrangement.
Use counters labelled with charge rather than colours alone. Place protons and neutrons centrally and electrons outside, then add the signed charges. Identify the element from proton number using the periodic table. Explain neutrality as cancellation, rather than saying the atom contains no charged particles. Model distances are deliberately not to scale.
Atoms of the same element with different numbers of neutrons.
In nuclide notation the upper-left number is A and the lower-left is Z. Protons=Z; neutrons=A−Z. A neutral atom has Z electrons; a positive ion has lost electrons and a negative ion has gained them. Chlorine-35 and chlorine-37 both have Z=17, with 18 and 20 neutrons respectively. Their mass numbers differ without a change in proton number.
Write a table with separate columns for A, Z, proton, neutron, electron and charge. Check A=protons+neutrons and charge=protons−electrons. For size comparisons convert radii into the same unit before dividing. A nucleus drawn clearly inside a circle is greatly enlarged relative to the atom; label the diagram not to scale and use numbers for a scale calculation.
The abundance-weighted average atomic mass relative to one twelfth of carbon-12.
For a model chlorine sample containing 75% chlorine-35 and 25% chlorine-37, Ar=(35×75+37×25)/100=35.5. The mean is closer to 35 because that isotope is more abundant. An ordinary unweighted mean would give 36 and incorrectly treat both abundances as equal. Ar is a relative quantity without a gram unit; it is not the mass number of an individual atom.
Check that percentage abundances total 100 before using the percentage formula. A missing second abundance can be found as 100 minus the first only when the sample has exactly two stated isotopes. Estimate where the answer should lie between the isotope masses, then calculate without premature rounding. A bead model with labelled masses illustrates weighted counting without implying all natural samples have the chosen fictional proportions.
An energy level occupied by electrons outside the nucleus in the atomic model.
The first twenty arrangements are H 1; He 2; Li 2,1; Be 2,2; B 2,3; C 2,4; N 2,5; O 2,6; F 2,7; Ne 2,8; Na 2,8,1; Mg 2,8,2; Al 2,8,3; Si 2,8,4; P 2,8,5; S 2,8,6; Cl 2,8,7; Ar 2,8,8; K 2,8,8,1; Ca 2,8,8,2. Verify every sum against the neutral atom’s proton number. This simplified sequence should not be extrapolated to every later element or used to claim the third shell can never contain more than eight.
Start with the atomic number from the supplied periodic table. For a neutral atom count that many electrons, fill the inner levels first and write a comma-separated arrangement. Convert it to a diagram by placing the exact number of dots on each ring and label the element. For an ion, adjust electron number before drawing and retain the unchanged nucleus; ion notation and brackets belong with the diagram.
A vertical column of elements with related chemical properties in the periodic table.
Group 1 atoms tend to lose one outer electron and form +1 ions; Group 7 atoms can gain one to form −1 ions. Predict similar reaction types within a group, then use the group-specific trend for relative reactivity. Helium has two outer electrons in a full first shell and belongs with Group 0, not Group 2. Hydrogen is an unusual non-metal and should not be described as an alkali metal simply from its printed position.
Use the supplied periodic table and numerical structures rather than memorizing an outline alone. Find Z, write the arrangement, count occupied shells, identify the outer count and compare the known group. State a prediction and the evidence for it. These main-group rules do not supply a complete electron-configuration treatment of transition metals.
A stated expected outcome that can be compared with new evidence.
Later discoveries filled gaps with elements whose properties agreed with the predictions, supporting the classification. Modern order follows atomic number rather than relative atomic mass. Isotopes explain why average mass need not rise strictly with proton number: different isotope masses and abundances influence the mean. The historical improvement was evidence-based prediction, not knowledge of electron shells that had yet to be discovered.
Compare a supplied early table with a modern periodic table. Identify a gap or reversed pair, write the property pattern and decide what discovery would support or challenge the prediction. Use fictional property values for numerical interpolation and label them as such. Agreement with one predicted property is useful evidence but not proof that every aspect of a model is correct.
An element characteristically conducting electricity and forming positive ions in its specified chemical reactions.
Sodium, 2,8,1, loses its outer electron to form Na⁺. Chlorine, 2,8,7, gains one to form Cl⁻. Oppositely charged ions can form a compound. Non-metal atoms also share electrons in covalent molecules, rather than always becoming isolated negative ions. Characteristic property patterns have exceptions: graphite conducts electricity and mercury is a liquid metal. Use more than one line of evidence.
Compare teacher-provided property cards, conductivity observations and electronic structures. Choose a classification, explain the electron change in a specified reaction and name a limitation of a physical-property inference. Use pre-approved specimens and supervised low-voltage conductivity apparatus; do not attempt reactions with unknown materials to discover their identity.
A Group 0 element whose atoms have stable outer-electron arrangements.
Neon is 2,8 and argon 2,8,8. Compare these stable structures with an alkali metal’s one outer electron or a halogen’s seven. Unreactivity and boiling point are different properties: a rising boiling point does not imply rising chemical reactivity. Given trend data, predict an ordering; a trend alone rarely determines an exact numerical value for an unmeasured element.
Use supplied structures and a labelled boiling-point table. Keep negative temperature ordering explicit: −186 °C is higher than −246 °C. Plot values when given, then distinguish interpolation from extrapolation. Atomic drawings are not evidence that helium needs another six electrons. More detailed intermolecular-force explanation is developed in Bonding, 4.2.
A Group 1 metal with one electron in its outer shell.
With oxygen, these metals tarnish and burn to form oxygen-containing solids; lithium burns crimson, sodium yellow and potassium lilac. The simple oxide model uses 4Li + O₂ → 2Li₂O; sodium and potassium burning products can include peroxide or superoxide, so do not invent a single oxide formula for every demonstration. Reactivity increases Li→Na→K because the outer electron is farther from the nucleus and more shielded by inner shells, so it is lost more easily despite increased nuclear charge.
Interpret teacher demonstrations or supplied observations under school risk assessment. Alkali-metal water reactions need approved very small quantities, protective screens and trained handling; students should not independently scale them up. Compare the same reaction conditions and distinguish qualitative vigour from a measured rate. Oxygen and chlorine demonstrations require their own approved controls.
A Group 7 non-metal element with seven outer electrons.
A halogen gains an electron to form a −1 halide ion when reacting with a metal, giving an ionic salt such as NaCl. With non-metals it generally forms covalent compounds through shared electron pairs, such as HCl. Down the group the outer shell is farther from the nucleus and more shielded, reducing attraction for an incoming electron despite greater nuclear charge. This is the opposite reactivity trend to Group 1, which loses an electron.
Compare supplied molecular drawings, states and reaction data without handling halogens independently. Chlorine and bromine require teacher-approved containment and ventilation; appearance cards can provide the required evidence safely. Use the same conditions when comparing reactivity. From a trend predict whether an unfamiliar halogen is more or less reactive, and avoid fabricating an exact boiling point.
A reaction in which a more reactive element replaces a less reactive element in a compound.
A displacement table should identify the added halogen, starting halide and expected product before recording colour. Bromine in water is orange/brown and iodine solutions can appear brown; observation must be compared with suitable blanks, because the added halogen also has colour. Same-halogen/halide combinations have no net displacement. A missing colour change alone may also reflect low concentration or observation limits.
For AT6 school-supervised work, use approved dilute halogen solutions in small-scale containers under the teacher’s risk assessment. Keep solution volumes and concentrations comparable, label each well and include reference colours. Do not generate chlorine or use an unapproved extraction solvent. Record actual observations, then compare with the predicted ranking; this activity is not one of the separately numbered required practicals.
A metal from the central transition region exemplified here by Cr, Mn, Fe, Co, Ni and Cu.
Iron can react with steam under suitable conditions and can rust slowly in oxygen and water; copper does not react with cold water in the Group 1 manner. Many transition metals react with oxygen or halogens when heated, so less reactive does not mean incapable of reaction. Compare the same conditions and specify whether an observation is rapid burning, slow corrosion or no visible change.
Use supplied property data and teacher-approved solid samples to compare density and qualitative hardness. Avoid student handling of alkali metals. Rank measured values rather than extrapolating an exact value for every transition element. For reaction comparisons, teacher demonstrations or source observations must have stated conditions and comparable sample sizes.
A substance increasing reaction rate by a lower-activation pathway without being used up overall.
Catalysts increase reaction rate without being used up overall. Iron catalyses the Haber process; nickel is used in hydrogenation; manganese dioxide catalyses hydrogen peroxide decomposition. A catalyst changes the route and rate, not the balanced reaction’s overall atom count. Variable charges require matching negative-ion charge: FeCl₂ contains Fe²⁺ and FeCl₃ contains Fe³⁺ with Cl⁻ in each case.
Use labelled compound photographs, supplied solution observations and given ion charges. Compare an actual teacher-supervised catalysed reaction with an uncatalysed control, maintaining temperature, concentration and volume. Cobalt/nickel compounds and strong oxidizers require approved handling and disposal; reference evidence is appropriate when they are not used in a school experiment. Do not identify an unknown compound solely from colour.
Use names and symbols of the first twenty elements, Groups 1 and 7, and other specified elements.
Distinguish atoms, elements and compounds; interpret fixed proportions and name compounds from formulae.
Write word equations, formulae and balanced symbol equations for specified reactions.
Explain why coefficients can change but the chemical formulae of substances cannot.
Write balanced half equations and ionic equations where appropriate.
Balance both atoms and charge; cancel spectator ions only when present unchanged on both sides.
Describe and explain filtration and crystallisation as physical separations.
Choose a sequence for mixtures containing insoluble solids and soluble substances.
Describe simple and fractional distillation and chromatography, with examples and explanations.
Select a suitable separation using solubility, boiling-point and movement evidence.
Explain why alpha-scattering evidence replaced the plum pudding model with the nuclear model.
Describe electron discovery, Bohr energy levels, protons and Chadwick neutron evidence without unrequired experimental detail.
State relative charges of proton, neutron and electron and use the nuclear model.
Explain neutral atoms and identify elements from atomic number.
Relate typical atom and nucleus radii using nano and standard form; locate almost all mass in the nucleus.
Use mass and atomic numbers to calculate protons, neutrons and electrons in atoms and ions.
Calculate relative atomic mass from given percentage isotope abundances.
Distinguish an abundance-weighted mean from an isotope’s integer mass number.
Represent the first twenty neutral atoms by numerical electronic structures and shell diagrams.
Place electrons in the lowest available energy levels using the specified GCSE model.
Relate position in the modern periodic table to atomic number, occupied shells and outer electrons.
Use group relationships to predict reactions and probable reactivity.
Describe early atomic-weight ordering, Mendeleev gaps and property-based reordering.
Explain how later discoveries and isotope knowledge supported development of the modern table.
Compare characteristic physical and chemical properties and relate them to periodic position.
Explain ion formation using outer electrons without treating all characteristic properties as universal.
Explain low reactivity and monatomic nature using stable outer-shell arrangements.
Predict boiling-point relationships from the increase down Group 0.
Describe Li, Na and K reactions with oxygen, chlorine and water.
Explain characteristic +1 ion formation and increasing reactivity down the group.
Describe diatomic halogens and the nature of compounds formed with metals and non-metals.
Explain decreasing reactivity and increasing molecular mass/melting/boiling points down Group 7.
Predict displacement of a less reactive halogen from an aqueous halide solution.
Use reaction observations and balanced equations to compare chlorine, bromine and iodine.
Compare melting point, density, strength, hardness and reactivity with oxygen, water and halogens.
Exemplify general differences with chromium, manganese, iron, cobalt, nickel and copper.
Exemplify different ion charges, coloured compounds and catalytic uses of Cr, Mn, Fe, Co, Ni and Cu.
Use supplied ion charges and conditions rather than assuming all compounds have one colour or charge.

element · ⁨元素⁩: A substance containing only one type of atom defined by proton number
compound · ⁨化合物⁩: A substance containing two or more elements chemically combined in fixed proportions
coefficient: A number before a chemical formula multiplying the entire formula
reactant · ⁨反应物⁩: A substance used up or changed in a chemical reaction
half equation: An equation showing electron loss or gain for one oxidation or reduction process
spectator ion: An ion unchanged on both sides of a reaction and omitted from its net ionic equation
residue: The insoluble solid retained by a filter
filtrate: The liquid or solution that passes through a filter
distillate: Liquid collected after vapour is condensed during distillation
chromatography: A separation method based on different distributions between stationary and mobile phases
nuclear model: An atomic model with mass and positive charge concentrated in a tiny central nucleus
alpha particle: A positively charged particle consisting of two protons and two neutrons
atomic number: The number of protons in an atom’s nucleus
ion · ⁨离子⁩: An atom or group of atoms carrying a net electric charge
isotope: Atoms of the same element with different numbers of neutrons
mass number: The total number of protons and neutrons in a nucleus
relative atomic mass: The abundance-weighted average atomic mass relative to one twelfth of carbon-12
isotope abundance: The proportion of atoms belonging to a particular isotope in a sample
electron shell: An energy level occupied by electrons outside the nucleus in the atomic model
electronic structure: The distribution of electrons among an atom’s energy levels
group: A vertical column of elements with related chemical properties in the periodic table
period · ⁨周期⁩: A horizontal row of elements in the periodic table
prediction: A stated expected outcome that can be compared with new evidence
atomic weight: The historical term used for relative atomic mass in early periodic classifications
metal · ⁨金属⁩: An element characteristically conducting electricity and forming positive ions in its specified chemical reactions
non-metal: An element generally lacking metallic properties, often gaining or sharing electrons in specified reactions
noble gas · ⁨稀有气体⁩: A Group 0 element whose atoms have stable outer-electron arrangements
boiling point: The temperature at which a liquid boils at a specified pressure
alkali metal: A Group 1 metal with one electron in its outer shell
shielding: Reduction in the attraction experienced by outer electrons due to inner electrons
halogen: A Group 7 non-metal element with seven outer electrons
diatomic molecule: A molecule containing two atoms
displacement reaction: A reaction in which a more reactive element replaces a less reactive element in a compound
halide ion: A negatively charged ion formed when a halogen atom gains an electron
transition metal: A metal from the central transition region exemplified here by Cr, Mn, Fe, Co, Ni and Cu
density · ⁨密度⁩: Mass per unit volume of a substance
catalyst · ⁨催化剂⁩: A substance increasing reaction rate without being used up overall
variable charge: The ability of an element to form ions with different charges in different compounds

4.2 · Bonding, structure and properties

Attraction between oppositely charged ions.
To explain a bulk property, name the structure, particles, forces and mobile charge carriers. Simple molecular substances can have strong covalent bonds inside molecules but weak attractions between molecules.
Compare substances using evidence such as melting point, conductivity when solid and molten, and solubility. One property rarely proves a structure; use a pattern of evidence.
Strong electrostatic attraction between oppositely charged ions.
Most non-metal elements and compounds of non-metals have covalent bonding, but their structures can be small molecules, very large molecules or giant networks. An ionic bond is the attraction after ion formation, not the electron transfer itself. A covalent pair is shared, not transferred completely to one nucleus. Metal electrons are free to move through the structure rather than assigned to one pair of atoms.
Sort teacher-provided particle diagrams by charges, shared pairs and delocalised electrons. For each classification state the particles, electrostatic attraction and whether the drawing represents a small molecule or part of a giant structure. Do not classify only from melting point: different structures can produce overlapping property data.
An electron diagram using different marks to show the origins of electrons.
Draw final ions in separate square brackets, with charges outside. A dot-and-cross diagram uses different marks for electrons originating from different atoms; transferred electrons keep their original marks. Mg²⁺ has no electrons left in its original third shell, while O²⁻ has six original outer electrons plus two transferred ones. If only outer electrons are shown, state whether the remaining filled shell of the cation is omitted by that convention.
Start from neutral structures and account for every transferred electron. For NaCl transfer one; for MgO transfer two; for MgCl₂ transfer one to each of two chlorine atoms; for Na₂O transfer one from each of two sodium atoms to oxygen. Confirm total charge zero and outer-shell completion. Dots and crosses indicate origin only: all electrons have the same physical charge and properties.
A repeating giant arrangement of ions held by electrostatic attraction.
To infer an empirical formula, count ions in a stated representative region and reduce the ratio to the smallest whole numbers. A model with six Mg²⁺ and twelve Cl⁻ gives MgCl₂ after reducing 6:12 to 1:2. Counts in an arbitrary boundary fragment need not themselves give the bulk ratio, so use the representative information or account for shared positions as specified. Knowledge of other named ionic crystal structures is not required.
Compare dot-and-cross, ball-and-stick and space-filling representations. Dot/cross shows electron origin and charge but not the full lattice geometry. Ball/stick makes an arrangement clear but rods are not physical ionic bonds and spaces/radii may be misleading. A 2D slice omits neighbours above and below; a 3D model can hide ions behind others. State the particular useful feature and limitation of each model.
A bond formed by a shared pair of electrons attracted to both nuclei.
Oxygen has six outer electrons, so O₂ has two shared pairs, a double bond, and two lone pairs on each atom. Nitrogen has five outer electrons, so N₂ has three shared pairs, a triple bond, and one lone pair on each atom. A line diagram writes H–H, Cl–Cl, H–Cl, O=O or N≡N. One line represents one shared pair, not one electron.
Count the available outer electrons first. Place shared pairs between the correct atoms using dots from one and crosses from the other, then place remaining electrons as lone pairs. Check each atom’s full outer shell and the total electrons drawn. Compare with a line or ball-and-stick model: these make connectivity clear but normally omit lone pairs and electron origin.
A substance containing very large molecules built from repeated units.
Covalent substances can consist of small molecules, very large molecules such as polymers, or giant networks such as diamond and silicon dioxide. A polymer has many repeated units linked along each molecule: [–CH₂–CH₂–]ₙ represents poly(ethene), with bonds passing through brackets and n large. A giant network diagram shows only a small part of continuing covalent connections, not a separate molecule with that exact pictured atom count.
For electron diagrams, use one mark type for the central atom’s electrons and another for hydrogen’s. Count bonds, lone pairs and all outer electrons. For a molecular formula, count each element in the complete stated molecule. For a polymer, identify the bracketed repeat rather than calling the drawing a small molecule. Ball-and-stick drawings enlarge gaps and do not fix scale; a 2D line formula omits the real 3D arrangement.
An electron not attached to one atom or bond and able to move through the structure.
The shared electrons extend throughout the structure, so shifting layers can preserve attraction and allow bending. When a potential difference is applied, mobile electrons carry charge; the positive ions remain in their positions in a solid rather than drifting through the wire. The drawing’s ions and electron symbols are a model of the bulk bonding, not isolated metal ions mixed with a separate substance.
Compare a metallic diagram with an ionic lattice and a small covalent molecule. Identify whether electrons are local shared pairs or delocalised and whether both positive and negative ions are represented. Use school-approved low-voltage conductivity demonstrations, with no mains circuit and no heating of unknown samples. A schematic can show carriers but not their detailed motions or actual relative sizes.
The change of state from solid to liquid.
Stronger relevant forces require more energy to overcome and commonly give higher melting and boiling points. Identify the particles and forces before explaining: boiling a molecular liquid overcomes intermolecular attractions, while melting an ionic lattice overcomes strong ion attractions. Bulk properties belong to large collections, not an individual atom that is itself a tiny piece of solid or liquid. At a transition temperature more than one state can coexist.
Given melting and boiling points at one pressure, compare the temperature with both boundaries. Below melting: solid; between: liquid; above boiling: gas. Use teacher-provided observations or an approved water-heating/cooling experiment. Record temperature and state, and distinguish a measured plateau from a schematic curve. Changing pressure or mixtures can alter transition temperatures.
A representation of matter using particles to explain arrangement and behaviour.
Without attractions, the drawing cannot explain energy needed to separate particles or the different melting and boiling temperatures of substances. Treat the balls as symbols for particles, not literal hard miniature pieces of bulk material. A model may succeed at showing close liquid particles while needing an extended force account to explain boiling. Individual particles do not acquire the substance’s bulk hardness or melting point.
Compare the same state-change story using a sphere diagram and a labelled force model. For every criticism state what is omitted and the consequence for the explanation. Keep useful features such as particle-number conservation. A model improvement is a justified new representation, not simply adding many details with no explanatory purpose.
A solution in which water is the solvent.
For an aqueous precipitation reaction, AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq) distinguishes the new solid from dissolved substances. Magnesium reacting with acid can be written Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g). Correct state labels add evidence about the reaction while the coefficients still conserve atoms.
Read the stated conditions and observation before labelling: a collected gas, a precipitate, a pure liquid or a solution. Do not infer that every formula containing H is gaseous or that every liquid mixture should be labelled (l). Compare balanced equations with observations from a teacher-approved demonstration and state what each label tells the reader.
An ion able to move through a liquid or solution and carry charge.
The carrier is an ion, not a free electron travelling through the salt. Dissolution separates ions into the solution; it does not simply melt the salt. Not every ionic compound is readily soluble, so explain aqueous conductivity only when a dissolved sample is specified. High melting point supports a strong giant structure, while state-dependent conductivity strengthens the ionic interpretation.
Compare teacher-provided data for solid, molten and dissolved samples. High-temperature molten-salt demonstrations require approved equipment and supervision; students can analyse supplied observations instead. For low-voltage solution tests, keep concentration, electrode spacing and immersion depth comparable and use clean apparatus. A brightness comparison is qualitative unless current is actually measured.
An attraction between separate molecules.
For comparable molecules, intermolecular forces generally increase with molecular size, so larger molecules tend to have higher melting and boiling points. Molecular shape and other interactions also matter; a size trend is not an exact universal temperature rule. A compound such as hydrogen chloride can form ions on dissolving, so the pure molecular substance and its reacting aqueous solution must be distinguished.
Use paired molecular drawings that label strong internal bonds and weaker attractions between separate molecules. Identify what is overcome during a specified physical change and what remains intact. Compare supplied boiling-point data for a related series at the same pressure. Do not infer a substance’s identity from boiling point alone without composition and other evidence.
A long polymer molecule containing many repeated units.
Distinguish bonds along a chain from attractions between separate chains. The size of polymer molecules gives many opportunities for intermolecular attraction. A simplified chain picture helps explain the solid material, but real polymer properties also depend on chain arrangement and other structural features. Do not replace intermolecular attractions with a claim that every neighbouring chain is necessarily joined by covalent bonds.
Compare a small molecular diagram, several long chain fragments and a covalent network. Identify repeated units, chain boundaries and the type of connection shown. Use prepared polymer samples or teacher-provided images, not unapproved heating or burning tests. The [–CH₂–CH₂–]ₙ model shows poly(ethene); the two bonds through the brackets show continuation of the chain.
An extended network of atoms connected by strong covalent bonds.
In silicon dioxide, each silicon is linked to four oxygens and each oxygen bridges two silicons, producing the overall Si:O ratio 1:2. A local drawing may show four surrounding oxygens because each is shared with another silicon elsewhere in the network. Graphite has strongly bonded layers but weaker attractions between layers; its high-temperature behaviour and layer sliding concern different connections.
Trace bonds from one atom to its neighbours and then onward beyond the fragment. Distinguish an uninterrupted network from separate groups with gaps between them. Compare a local coordination drawing and the bulk formula. Models may omit remote bonds, distort angles or enlarge spaces; explain what the representation does and does not establish.
A metallic mixture containing a metal and one or more other elements.
Different-sized atoms distort the regular layers and make sliding more difficult, so many alloys are harder than the corresponding pure metal. Hardness is resistance to indentation or scratching, while strength concerns resisting deformation or failure under a load; do not treat the words as identical measurements. An alloy still has metallic bonding and need not have one fixed compound formula.
Compare equal-size atom layers with a model containing some larger atoms. Show the disruption and explain its effect rather than merely saying alloy atoms are bigger. Use supplied hardness data and actual teacher-approved samples with suitable tools. Compare the same test method and conditions; an arbitrary harder sample does not establish how every possible alloy behaves.
Transfer of thermal energy through a material from hotter toward cooler regions.
Electrical conduction is transfer of charge; thermal conduction is transfer of energy from a hotter region toward a cooler region. Do not confuse electron drift with the whole wire moving or with electrons being permanently consumed. A metal can conduct thermal energy without being part of an electric circuit. The two explanations share mobile electrons but refer to different observations.
For a supervised low-voltage electrical comparison, keep geometry and temperature controlled and measure voltage and current if interpreting quantitative data. For heat transfer use teacher-approved apparatus and handling, because hot metal may look unchanged. Analyse supplied measurements if direct heating is unsuitable. The simple carrier diagram does not show the detailed speed or microscopic energy distribution.
A carbon form with each atom covalently bonded to four others in a giant three-dimensional structure.
Explain one property with the appropriate connection: resistance to deformation follows the rigid many-direction network; high melting point follows energy needed to overcome many strong covalent bonds; non-conductivity follows the absence of mobile charged carriers. A drawing with a carbon surrounded by four neighbours shows local coordination, not a complete diamond molecule.
Compare a tetrahedral local model with a graphite-layer model. Count four bonds around an interior diamond carbon and trace continuation beyond the fragment. Mark a two-dimensional picture as a projection of a three-dimensional arrangement. Prepared models or images suffice; there is no need to heat or scratch expensive specimens or infer properties from a gem’s colour.
A carbon form with covalently bonded hexagonal layers and one delocalised electron per carbon.
Do not use weak interlayer forces to explain the high melting point: that property involves strong covalent bonding. Conversely, naming strong covalent bonds alone does not explain easy layer sliding. Graphite resembles metals in having delocalised electrons, but its bonding network is layered covalent rather than metallic. Conductivity directions in real graphite are not represented fully by one simple flat drawing.
Inspect a hexagonal-layer model and identify three neighbours for an interior carbon, with bonds continuing at cut edges. Compare layers stacked without covalent connecting lines. A teacher-approved low-voltage test of a prepared graphite rod can provide conductivity evidence. Record contact quality and geometry, and distinguish the rod material from a pencil’s mixed graphite/clay core.
A single carbon layer with the hexagonal bonded arrangement of graphite.
A composite combines materials so their useful properties can contribute to a designed product. Graphene can reinforce a material, but strength claims need actual data for the particular composite, loading and manufacturing method. Conductivity supports electronic applications without proving every graphene-containing product is automatically a good conductor. A drawn lattice edge is a clipped boundary, not the full bulk bonding environment.
Compare a single honeycomb sheet, several graphite layers and a hollow fullerene shape. Identify graphene from the single-layer description and explain one property with bonding or electrons. Use supplied product data rather than asking students to handle loose nanopowders or claiming laboratory-quality graphene has been produced from an ordinary pencil drawing.
A molecule of carbon atoms arranged in a hollow cage or related hollow shape.
A hollow molecular cage can be considered for carrying other substances in proposed delivery systems; nanotubes can be used as reinforcing fibres and in electronic materials. A use must be linked to supplied properties and tested for suitability. Do not claim every nanotube has identical conductivity or that a proposed medical use is automatically safe or effective. Graphene is a sheet, graphite stacked layers and a fullerene a hollow shape.
Compare labelled structural descriptions with shape drawings. Distinguish a schematic outline showing geometry from a complete atom-and-bond map: a circle alone does not establish exactly sixty carbons. For nanotube data convert length and diameter into the same unit before finding their ratio. Use contained prepared materials or reference images rather than dispersing loose nanomaterials.
A particle with dimensions in the approximate 1–100 nm range used in this specification.
For a cube side L, area=6L², volume=L³ and area/volume=6/L. Reducing side length tenfold raises area/volume tenfold. Nanoparticles can have different properties from the same bulk material because a much larger fraction interacts at the surface. More exposed area can improve catalytic effectiveness per mass, but particle shape, aggregation and surface chemistry also affect performance.
Use cubical models as a mathematical illustration and keep length units consistent. Area/volume is a ratio with reciprocal-length units, not simply a percentage. Compare a nano dimension with a stated atomic size by division. A cube model does not imply real nanoparticles are all cubes or all have the same chemical behaviour.
Use of nanoscale structures and their properties in designed applications.
Consider useful performance, required quantity, cost and durability alongside possible human exposure, environmental release and uncertain long-term effects. Risk depends on the material, particle form, exposure route and amount. A contained particle in a product and an inhalable loose powder need separate exposure evidence. A proposed benefit is not proof of safety, while a possible risk is not proof that harm occurs at every dose.
Use a supplied product evidence table, identifying measured outcomes, comparison conditions and missing information. Weigh advantages and disadvantages for the intended function and make a qualified recommendation supported by evidence. Do not ask students to make a cosmetic, administer a nanomaterial or generate nanopowder. Any actual school materials require the approved method and containment.
Explain ionic, covalent and metallic bonding using electrostatic forces and electron transfer or sharing.
Identify typical metal/non-metal combinations without confusing bond type with bulk structure.
Draw outer-electron diagrams for Group 1/2 metals with Group 6/7 non-metals.
Infer ion charges and noble-gas electronic structures from these group numbers.
Recognise ionic structures and explain strong attractions in all directions.
Deduce empirical formulae from representative ion ratios and evaluate dot/cross, ball/stick and 2D/3D models.
Draw dot-and-cross diagrams for H₂, Cl₂, O₂, N₂ and HCl.
Represent single, double and triple bonds and retain the correct unshared electrons.
Draw electron diagrams for H₂O, NH₃ and CH₄ and deduce molecular formulae from models.
Represent polymer repeat units and parts of giant covalent structures; evaluate model limitations.
Recognise and explain metallic giant structures with delocalised outer electrons.
Distinguish metallic attraction from covalent electron pairs and ionic-compound lattices.
Explain melting, freezing, boiling and condensation using particles, energy and forces.
Predict states from melting/boiling data and distinguish bulk properties from individual-atom properties.
Explain limitations of solid, inelastic sphere models with no represented forces.
Evaluate the model in relation to changes of state without discarding its useful arrangement comparison.
Use (s), (l), (g) and (aq) correctly in specified reaction equations.
Distinguish a dissolved solute from its pure solid, liquid or gas form.
Explain high melting/boiling points through many strong electrostatic attractions.
Explain conductivity when molten or dissolved and absence of ion mobility in a solid.
Explain low melting/boiling points and non-conductivity of typical small molecular substances.
Relate increasing molecular size to stronger intermolecular forces in comparable substances.
Recognise polymers from bonding/structure diagrams.
Explain strong covalent bonds within polymer chains and relatively strong attractions between large molecules.
Recognise giant covalent structures from diagrams and distinguish them from small molecules and polymers.
Explain very high melting points using many strong covalent bonds; include diamond, graphite and silicon dioxide.
Explain metal melting/boiling points and malleability using strong metallic bonding and layers.
Explain why mixing different atoms can make an alloy harder than its pure metal.
Explain electrical conductivity through delocalised-electron movement.
Explain thermal conduction through energy transfer by delocalised electrons.
Explain hardness and very high melting point through four covalent bonds per carbon.
Explain absence of electrical conductivity because no delocalised electrons are available.
Explain three covalent bonds per carbon, hexagonal layers and no covalent bonds between layers.
Explain sliding, high melting point and conductivity with the correct parts of the structure.
Recognise graphene as a single graphite layer and explain its properties from bonding.
Connect strength and electrical conductivity to potential composite/electronic uses.
Recognise hollow fullerene molecules, C₆₀ and cylindrical nanotubes from descriptions/diagrams.
Give uses linked to structure, including nanotechnology, electronics and materials.
Compare nanoparticles with atom/molecule, fine-particle and coarse-particle dimensions using SI prefixes.
Calculate cube surface/volume and explain how shrinking size changes the ratio and possible properties.
Evaluate supplied advantages/disadvantages for applications in medicine, electronics, cosmetics, deodorants or catalysis.
Explain possible risks and evidence needed without inventing unrequired specific material properties.

ionic bond · ⁨离子键⁩: Attraction between oppositely charged ions
delocalized electron: An electron not confined to one atom or bond
ionic bonding: Strong electrostatic attraction between oppositely charged ions
metallic bonding: Electrostatic attraction between positive metal ions and delocalised electrons
dot-and-cross diagram: An electron diagram using different marks to show the origins of electrons
electron transfer: Movement of electrons from one atom to another when ions form
ionic lattice: A repeating giant arrangement of ions held by electrostatic attraction
empirical formula: The simplest whole-number ratio of atoms or ions in a compound
covalent bond · ⁨共价键⁩: A bond formed by a shared pair of electrons attracted to both nuclei
lone pair: A pair of outer electrons not shared in a covalent bond
polymer: A substance containing very large molecules built from repeated units
repeat unit: The group of atoms repeated along a polymer molecule
delocalised electron · ⁨离域电子⁩: An electron not attached to one atom or bond and able to move through the structure
electrostatic attraction · ⁨静电吸引力⁩: Attraction between opposite electrical charges
melting · ⁨熔化⁩: The change of state from solid to liquid
condensation: The change of state from gas to liquid
particle model: A representation of matter using particles to explain arrangement and behaviour
model limitation: A feature omitted or simplified that restricts a model’s explanatory use
aqueous solution: A solution in which water is the solvent
state symbol: A label specifying a substance’s physical state or aqueous form in an equation
mobile ion: An ion able to move through a liquid or solution and carry charge
electrical conductivity: The ability of a material to carry electrical charge through itself
intermolecular force: An attraction between separate molecules
small molecule: A discrete covalently bonded group containing relatively few atoms
polymer chain: A long polymer molecule containing many repeated units
covalent backbone: The sequence of atoms joined by covalent bonds along a polymer molecule
giant covalent structure: An extended network of atoms connected by strong covalent bonds
silicon dioxide: A substance with a giant covalent structure and overall silicon-to-oxygen ratio 1:2
alloy: A metallic mixture containing a metal and one or more other elements
hardness: Resistance of a material to indentation or scratching
thermal conduction: Transfer of thermal energy through a material from hotter toward cooler regions
charge carrier: A mobile charged particle that contributes to an electric current
diamond · ⁨金刚石⁩: A carbon form with each atom covalently bonded to four others in a giant three-dimensional structure
rigid network: An extended arrangement whose strong connections resist changes in shape
graphite: A carbon form with covalently bonded hexagonal layers and one delocalised electron per carbon
hexagonal ring: A ring of six connected atoms in the stated graphite layer model
graphene: A single carbon layer with the hexagonal bonded arrangement of graphite
composite: A material combining components to obtain a useful combination of properties
fullerene: A molecule of carbon atoms arranged in a hollow cage or related hollow shape
carbon nanotube: A cylindrical fullerene with a very large length-to-diameter ratio
nanoparticle: A particle with dimensions in the approximate 1–100 nm range used in this specification
surface-area-to-volume ratio: Surface area divided by volume, measuring exposed area relative to material volume
nanotechnology: Use of nanoscale structures and their properties in designed applications
exposure route: The way a person or environment comes into contact with a material

4.3 · Quantitative chemistry

The unit of amount containing 6.02×10²³ stated particles at the GCSE approximation.
Calculate the amount available for each reactant and divide by its coefficient. The smaller ratio limits the reaction. Use that reactant to calculate the maximum product before comparing actual yield.
Write the balanced equation first, include units in molar masses, then convert each given mass or solution volume into amount. Convert cubic centimetres to cubic decimetres before using concentration in moles per cubic decimetre.
The total mass of all reactants equals the total mass of all products in a chemical reaction.
For 2H₂ + O₂ → 2H₂O, the coefficient 2 multiplies an entire formula. The subscript 2 belongs to the preceding element, so one water molecule contains two H atoms and one O atom. Both sides contain four H atoms and two O atoms. Using supplied relative atomic masses H=1 and O=16 gives reactant totals 2×2 + 32 =36 and product total 2×18=36. The coefficients 2:1:2 are not gram ratios.
Make an atom-count table before doing a mass calculation. Identify the physical boundary of the weighed system: vessel, liquid, solid and any contained gas. For an actual supervised demonstration, use an approved arrangement that accommodates any pressure change; do not improvise a heated sealed vessel. Compare readings only after correcting for anything added or removed from the balance.
The sum of relative atomic masses for the atoms shown in a formula.
An element’s percentage by mass is its total relative mass contribution divided by Mr, multiplied by 100. Calcium contributes 40 of the 74 units in Ca(OH)₂, giving 40/74×100=54.1% to three significant figures. Oxygen contributes 32 and hydrogen 2. These percentages describe the compound’s fixed mass proportions, not the proportion of different atoms counted equally.
List element, atom count, Ar and mass contribution in four columns. Add contributions for Mr, then select the requested element’s contribution for the numerator. In a balanced equation include coefficients: 2Mg + O₂ → 2MgO gives 2×24+32=80 on the left and 2×40=80 on the right. Use the periodic-table values supplied with a particular question rather than guessing more precise data.
A system that can exchange material with its surroundings.
Identify which material is on the balance at each reading. In the metal example, oxygen joins the weighed solid; in the carbonate example, gas leaves the weighed apparatus. In particle terms the same atoms are redistributed among substances, with some crossing the selected system boundary. A sealed-system total includes all products; an open-residue measurement does not.
For a teacher-approved school investigation, record the clean dry vessel mass, sample mass and final cooled vessel-plus-solid mass. Use appropriate heat protection, ventilation and an approved method. Loss of solid through spitting or incomplete reaction can change the reading too. Do not seal heated gas-producing apparatus merely to demonstrate conservation, and do not assume every measured difference is the intended gas transfer.
An estimate of the doubt associated with a measurement result.
For 2.40, 2.42 and 2.44 g, the mean is 2.42 g and the full range is 0.04 g. Half-range is 0.02 g, so the stated estimate is 2.42±0.02 g. Here the mean lies midway between the extremes. For an asymmetric distribution the extreme deviations from the mean differ; display the actual results as well as the chosen uncertainty estimate.
Plot a dot for each repeat on a labelled mass axis; coincident results can stack vertically. Keep units and sensible decimal places. Investigate an unusual value before excluding it, record any exclusion reason and preserve the original readings. More repeats reveal random spread but cannot by themselves remove a systematic bias, such as an uncorrected balance zero.
State what is being counted. A mole of CO₂ molecules contains one mole of carbon atoms and two moles of oxygen atoms, three moles of atoms altogether. A mole of NaCl formula units corresponds to one mole of Na⁺ ions and one mole of Cl⁻ ions in the ionic lattice. Calling either sample simply a mole of particles without specifying which particles can make an answer ambiguous.
Write mass, molar mass and amount in separate columns. Convert kilograms or milligrams into grams before dividing. Rearrange n=m/M to m=nM or M=m/n as needed. Calculate particle number as n×6.02×10²³ and retain standard form; divide particle number by the same constant to find amount. Use sensible significant figures and keep relative mass distinct from molar-mass units.
The ratio of reacting amounts given by balanced equation coefficients.
For 2Mg + O₂ → 2MgO, supplied molar masses Mg=24, O₂=32 and MgO=40 g per mol give mass proportions 48:32:80. These differ from coefficients 2:1:2. A calculation from one reactant assumes sufficient other reactant and complete reaction unless the question supplies a limiting amount. If the known quantity is product mass, work backwards through the same ratio.
Balance the equation first and underline the known and requested substances. Write a mole row underneath the coefficients, labelling every substance. Avoid rounding a small mole amount early. State the sufficient-reactant and complete-conversion assumptions. Numerical theoretical masses do not include recovery loss, side reactions or reversible-equilibrium effects.
A whole-number coefficient before a formula in a balanced equation.
For Mg, O₂ and MgO with molar masses 24, 32 and 40 g per mol, supplied masses 2.4, 1.6 and 4.0 g give 0.10, 0.050 and 0.10 mol. Dividing by 0.050 gives 2:1:2, hence 2Mg + O₂ → 2MgO. The oxygen formula is O₂, so its molar mass is 32, not the atomic value 16.
Organise substance, mass, molar mass, amount and simplified ratio in a table. When a ratio includes 1.5, multiply every term by two instead of rounding that term to one or two. Check the resulting atom counts and conserved total mass. Reacting masses must exclude unreacted excess or unrelated apparatus; otherwise they cannot directly establish the equation ratio.
The reactant used up first under the stated reaction model, limiting possible product.
For Mg + 2HCl → MgCl₂ + H₂, 0.20 mol Mg would require 0.40 mol HCl. If only 0.30 mol HCl is supplied, acid limits and only 0.15 mol Mg can react. Maximum H₂ amount is 0.15 mol, and 0.05 mol Mg remains. This assumes reaction proceeds as written and does not describe incomplete conversion caused by slow rate or equilibrium.
Write available amount, required coefficient and available amount divided by coefficient for each reactant. The smaller scaled supply determines the maximum reaction extent. Multiply that extent by the target product coefficient. To find excess remaining, subtract the amount consumed from the amount supplied. Avoid declaring a limiting reactant solely from a visual guess or whichever mass number is smaller.
The substance dissolved in a solvent to form a solution.
For a uniformly mixed solution, the amount of solute taken is proportional to the sample volume. A 0.250 dm³ sample of an 8.0 grams per dm³ solution contains 2.0 g solute. The unit calculation cancels dm³ and leaves grams. This shared-tier numerical task is distinct from the embedded Higher-only explanation of how changing mass or solution volume changes concentration.
Write the concentration unit beside the value and convert volume before multiplying. Mark whether a number refers to solute, solvent or solution. For an actual teacher-approved preparation, weigh the selected solute, dissolve it and make the solution up to the required final volume using appropriate school apparatus. Check complete dissolution and safe handling; the written calculation does not replace actual supervised technique.
Lowering concentration by adding solvent while retaining the stated solute amount.
During simple dilution with pure solvent and no losses or reaction, the solute mass stays constant. Therefore c₁V₁=c₂V₂ when both concentrations use the same mass units and both volumes the same volume units. The final volume includes the original sample and added solvent; it is not the volume of water added alone. Dissolution and dilution should not be confused with a chemical transformation of solute.
Identify what is held constant before stating a direction of change. Rearrange c=m/V to V=m/c or m=cV. For a planned school dilution, use appropriate measured transfer and make up to the required final volume under teacher supervision. Concentrated acids require the school’s specific safe procedure; the generic arithmetic is not an instruction to mix hazardous liquids.
The amount of desired product obtained.
Recovery may be below the maximum because a reversible reaction does not go to completion, product is lost during separation, or reactants undergo side reactions. Some atoms can remain in reactants or enter unwanted products; some desired product can remain dissolved or on apparatus. Conservation concerns all atoms and substances, while yield concerns the recovered desired product.
Record a dry product mass rather than including retained solvent or unrelated material. Write actual and theoretical values in the same unit before dividing. Inspect a proposed explanation against the method: transfer loss, incomplete crystallisation or reversible conversion are different causes and need different evidence. School-supervised separations remain actual tasks; this written evaluation supports them rather than replacing them.
The maximum desired product amount predicted by the stated reaction model.
For CaCO₃ → CaO + CO₂, the amount ratio CaCO₃:CaO is 1:1. Supplied molar masses are 100 and 56 g per mol. A 10.0 g CaCO₃ sample contains 0.100 mol and can give 0.100 mol CaO, mass 5.60 g. The other 4.40 g is theoretical carbon dioxide, so the product-residue mass is lower than the original carbonate mass even at complete conversion.
Label theoretical and actual masses clearly. State sample purity and which reactant limits if relevant. Use unrounded intermediate values when multiplying a ratio. Check the total product masses against reactant mass for a decomposition with no added reactant. For an actual school experiment, follow approved heating and handling procedures; calculation alone does not establish purity or completion.
The fraction of starting-material mass directed into the stated desired product by the balanced equation.
For CaCO₃ → CaO + CO₂ with relative masses 100, 56 and 44, choosing CaO as desired product gives 56/100×100=56%. Choosing CO₂ as the desired product for a different purpose gives 44%. Conservation still accounts for all 100 mass units. A high atom economy can reduce waste and raw-material cost, but it does not alone establish safety, energy demand or commercial suitability.
Balance the equation, identify the desired product explicitly, and write every coefficient-weighted relative mass. For 2H₂ + O₂ → 2H₂O, desired water contribution is 2×18=36 and reactant total is 2×2+32=36, so atom economy is 100%. This uses relative mass bookkeeping rather than an experimental product weighing. Both-tier calculations are distinct from Higher-only comparison of complete reaction pathways.
A chosen sequence or route of reactions for producing a substance.
Imagine route A with atom economy 90%, yield 50% and a slow rate, and route B with atom economy 70%, yield 90% and a faster rate under supplied comparable conditions. Neither percentage alone decides every use. If equal starting mass could theoretically enter product according to these fractions, recovered desired mass per 100 g starting material is 45 g for A and 63 g for B under this simplified comparison.
State what the data hold constant: feedstock mass, purity, conditions, reaction time or cost. Rank criteria against the question’s objective and acknowledge one disadvantage of the chosen route. Do not equate fast reaction with low energy cost unless the conditions support that inference. A favourable equilibrium position is different from a catalyst speeding the approach to it.
The amount in moles of stated solute per volume of solution.
Convert cm³ to dm³ by dividing by 1,000. A 250 cm³ portion of 0.20 moles per dm³ solution contains 0.20×0.250=0.050 mol solute. If solute molar mass is 40 g per mol, that amount weighs 2.0 g. At fixed amount, greater final solution volume means lower molar concentration. At fixed final volume, more solute moles means higher concentration.
Write units beside concentration and volume before multiplying. Identify which chemical species concentration describes. A bottle’s molar concentration normally refers to the stated solute formula amounts, not automatically the total amount of every ion after dissociation. Actual solution preparation requires a measured final volume, complete dissolution and the teacher-approved apparatus and procedure.
The volume delivered from a burette to reach the stated end-point.
A pipette supplies the measured sample volume and a burette delivers the titre. The titre is final minus initial burette reading, not the final reading alone. Use consistent close results under the school’s stated concordance rule and exclude a rough trial from the accurate mean. A suitable indicator end-point approximates the required neutralisation condition; it is not evidence that every solution must be pH 7.
Perform titration only as an actual supervised school task using approved dilute solutions, eye protection and the specified pipette filler. Rinse and use the apparatus as instructed, remove the filling funnel before readings and add dropwise near the end-point. Record readings at appropriate precision. These calculation exercises do not count as completing Required Practical 2, whose real method belongs in chemical changes.
The volume occupied by one mole of gas at specified conditions.
Equal gas amounts occupy equal volumes at identical conditions, so balanced coefficients give gaseous volume ratios. N₂(g) + 3H₂(g) → 2NH₃(g) gives ratio 1:3:2 for gas volumes measured at the same temperature and pressure. This ratio does not apply to a liquid or solid volume. A predicted product volume assumes sufficient reactants and the stated complete-conversion model; real equilibrium can reduce conversion.
Convert cm³ to dm³ consistently and label every volume’s conditions. Do not use 22.4 dm³ per mol, a value associated with different reference conditions, in this RTP task. Use teacher-provided gas data or approved low-risk gas collection under supervision; do not propose making ammonia or hydrogen in an improvised apparatus. Water-vapour and collection losses can affect real measurements.
Explain conservation of atoms and mass in chemical reactions.
Interpret equation coefficients and formula subscripts without confusing particle and mass ratios.
Calculate relative formula mass from atom counts and supplied relative atomic masses.
Calculate an element’s mass percentage in a compound and check mass totals in a balanced equation.
Explain apparent mass gain/loss using gas transfer and balanced equations.
Use the particle model to distinguish mass of the weighed residue from mass of all products.
Represent repeated-result distributions and estimate uncertainty.
Use the range around the mean without confusing uncertainty with a known error or instrument resolution.
Use n=m/Mr and m=nMr with mass in grams, and interpret molar mass correctly.
Use 6.02×10²³ particles per mole for stated atoms, molecules, ions or electrons.
Interpret balanced coefficients as mole ratios and calculate reactant/product masses.
Work from either a supplied reactant mass or a product mass without treating coefficients as gram ratios.
Convert supplied reactant/product masses to moles and find simple whole-number ratios.
Balance an equation from that ratio and check atoms; rearrange the amount equation when needed.
Explain limiting and excess reactants using balanced mole ratios.
Calculate the possible product from the limiting amount and any stated remaining excess.
Calculate solute mass from solution volume and mass concentration.
Use grams per dm³ consistently and distinguish solute mass, solution volume and solvent volume.
Explain c=m/V: increasing solute mass at fixed volume increases concentration.
Explain dilution at fixed solute mass and compare mass/volume changes quantitatively.
Calculate percentage yield when actual and theoretical product masses are supplied.
Explain incomplete reaction, recovery loss and side reactions while preserving conservation of atoms.
Calculate theoretical product mass from reactant mass and the balanced equation.
Combine that result with actual recovery to find yield, retaining complete-reaction and excess assumptions.
Calculate atom economy from desired-product and total-reactant relative masses, including coefficients.
Explain sustainability/economic value and distinguish atom economy from experimental yield.
Evaluate pathways using supplied atom economy, yield, rate, equilibrium and useful by-product evidence.
Justify a choice for stated priorities instead of selecting the largest single percentage automatically.
Calculate amount and solute mass from molar concentration and volume.
Explain how molar concentration relates to solute mass, molar mass and final solution volume.
Determine an unknown solution concentration from two reacting volumes and one concentration.
Apply balanced ratios and distinguish measured titre from pipetted sample and rough trial.
Use 24 dm³ per mol at stated RTP (20 °C, 1 atmosphere) to connect gas amount, volume and mass.
Use balanced ratios for gaseous reactant/product volumes only under the same temperature and pressure; rearrange formulae.

mole: The unit of amount containing 6.02×10²³ stated particles at the GCSE approximation
limiting reagent: The reactant that limits the possible product amount
conservation of mass: The total mass of all reactants equals the total mass of all products in a chemical reaction
system boundary: The defined limit separating the material being considered from its surroundings
relative formula mass: The sum of relative atomic masses for the atoms shown in a formula
mass percentage: An element’s mass contribution divided by the compound’s total mass, multiplied by 100
open system: A system that can exchange material with its surroundings
thermal decomposition: Breaking down a compound into simpler substances by heating
uncertainty: An estimate of the doubt associated with a measurement result
range · ⁨范围 (range)⁩: The largest result minus the smallest result in a set
Avogadro constant: The number of stated particles per mole, here 6.02×10²³ per mol
mole ratio: The ratio of reacting amounts given by balanced equation coefficients
stoichiometry: Quantitative relationships between substances in a balanced reaction
balancing number: A whole-number coefficient before a formula in a balanced equation
reacting mass: The mass actually consumed or formed in the stated reaction
limiting reactant: The reactant used up first under the stated reaction model, limiting possible product
excess reactant: A reactant supplied beyond the amount required to consume the limiting reactant
solute: The substance dissolved in a solvent to form a solution
mass concentration: The mass of solute per stated volume of solution
dilution: Lowering concentration by adding solvent while retaining the stated solute amount
final solution volume: The total volume of solution after the stated preparation or dilution
actual yield: The amount of desired product obtained
percentage yield: Actual yield divided by theoretical yield, multiplied by 100
theoretical yield: The maximum desired product amount predicted by the stated reaction model
complete conversion: The stated model in which the relevant limiting reactant is fully converted as written
atom economy: The fraction of starting-material mass directed into the stated desired product by the balanced equation
by-product: A product other than the stated desired product of a reaction
reaction pathway: A chosen sequence or route of reactions for producing a substance
process criterion: A stated factor used to compare production methods
molar concentration: The amount in moles of stated solute per volume of solution
aliquot: A measured portion of a larger sample of solution
titre: The volume delivered from a burette to reach the stated end-point
end-point: The observed change used to decide when to stop a titration
molar gas volume: The volume occupied by one mole of gas at specified conditions
room temperature and pressure: Here the specified reference conditions of 20 °C and 1 atmosphere

4.4 · Chemical changes

Loss of electrons.
Predict products using the specified electrolyte and electrode material. In an aqueous solution, hydrogen or oxygen may form because water-related species compete. Molten salts contain only the ions of the salt.
Use a low-voltage direct-current supply, approved electrodes, and the school risk assessment. Collect gases only by an approved method. Keep chlorine demonstrations teacher-controlled; do not ask students to generate hazardous gases independently.
The volume delivered from a burette to reach the stated end-point.
Calculate the known amount first, apply the stoichiometric ratio, then divide by the unknown solution volume in cubic decimetres. Use concordant titres as required by the school method and report the accepted values.
Rinse the burette with its solution and the pipette with the solution it transfers. Rinse the flask with distilled water. Add titrant slowly near the endpoint, swirl, and read the meniscus at eye level. Use a white tile and appropriate eye protection.
Oxidation identified as gain of oxygen.
For 2CuO + C → 2Cu + CO₂, copper oxide loses oxygen and is reduced, while carbon gains oxygen and is oxidised. The same oxygen atoms leave the oxide and appear in carbon dioxide. The two changes occur together in this reaction. Oxygen is an element; oxide is a compound containing oxygen with another element, so the names are not interchangeable.
Use atom models or supplied before/after formulae to trace oxygen. Balance each element and identify the substance receiving or losing oxygen. Actual burning-metal or heated-oxide work needs the teacher-approved method, heat protection and ventilation. A bright magnesium flame must not be viewed directly; a written example is not an instruction to ignite a sample independently.
An ordering of metals by their tendency to react and form positive ions.
Potassium, sodium and lithium react vigorously with cold water to produce hydroxides and hydrogen, with potassium generally most vigorous. Calcium also reacts, producing hydrogen and calcium hydroxide. Magnesium reacts very slowly with cold water, while zinc, iron and copper show no appreciable reaction under this school comparison. With dilute hydrochloric or sulfuric acid, Mg, Zn and Fe produce hydrogen, with Mg more vigorous and Fe slower; Cu does not. Very reactive alkali metals are not student acid-test samples.
Compare suitable teacher-approved metal samples under controlled acid concentration, volume, temperature and exposed surface conditions. Observe bubbles and metal consumption, but do not treat raw bubble counts from different-sized samples as a pure reactivity measure. Use supplied evidence or a teacher demonstration for highly reactive metals. Steam reactions and detailed unusual conditions are outside this stated recall scope.
A reaction in which a more reactive element replaces a less reactive element in a compound.
For supplied results, translate each successful displacement into a comparison arrow: A displaces B implies A>B. Combine comparisons into an order and check for contradictions. If A displaces B and B displaces C, A>B>C is consistent. Lack of visible reaction alone needs care because an unsuitable method, protective layer or short observation time can hide a reaction.
Use approved dilute salt solutions, clean matched metal samples, fixed volumes and comparable observation times. Record colour, deposit and sample changes without assigning identity from colour alone. A control with no added metal helps distinguish background change. Dispose of metal-containing solutions through the school’s specified collection route; do not test highly reactive metals in water-based salt solutions independently.
A metal occurring naturally in its elemental form.
For 2CuO + C → 2Cu + CO₂, copper oxide is reduced and carbon oxidised in oxygen terms. Carbon cannot be assumed to reduce every metal oxide under useful extraction conditions. Aluminium is more reactive than carbon and its extraction uses electrolysis, not simply this copper-oxide route. Detailed blast-furnace process recall is not required by this section’s stated limits.
When given an extraction description, locate the feed compound, reducing substance, desired metal and by-products. Evaluate supplied energy, cost, purity or waste data against a stated purpose rather than inventing industrial details. Any actual heated-oxide school work uses the approved small-scale method and technician-selected substances; industrial extraction is a reference case, not a school construction task.
Oxidation defined as loss of electrons.
Check both atom counts and total charge. The combined left charge is +2 and the right charge is +2. Zinc atoms become positive ions by losing electrons; copper ions become neutral metal atoms by accepting them. Identify the actual species reduced, Cu²⁺, rather than saying the already-neutral copper product gains further electrons.
Write the two half equations, multiply them if needed to make equal electron numbers, then cancel electrons. Cancel only unchanged spectator ions in a full ionic equation. Use the reactivity series to predict whether a proposed displacement is supported; a charge-balanced equation alone does not guarantee that the reaction will occur.
A suitable metal reacting with acid to produce a salt and hydrogen.
The metal must be suitable relative to hydrogen. Copper does not normally release hydrogen from these dilute acids. Do not extend this acid–metal pattern to nitric acid or concentrated acid reactions, which are outside the stated scope and can produce different outcomes. With a metal oxide, the usual products are salt and water instead of hydrogen.
Use teacher-approved dilute acid and selected small metal samples with eye protection and no nearby ignition source. Record bubbles and metal disappearance. A permitted small collected hydrogen sample gives a squeaky pop with the teacher-approved lit-splint test; do not apply a flame to the generating vessel or a large gas accumulation. Temperature, surface area and acid concentration affect the observed rate.
The H⁺ ion present in an aqueous acid in this GCSE model.
Charge is +2 on both sides of the net equation. Two H⁺ ions are needed to form one H₂ molecule; gaining one electron each makes their total gain two electrons. The metal product is an ion, while the hydrogen product is a neutral diatomic molecule. Electron gain must be assigned to hydrogen ions, not to chloride ions just because hydrochloric acid was used.
Separate the reactant species before identifying redox. In a supplied equation locate any metal atom becoming a positive ion and hydrogen ions becoming gas. Balance each half equation by atoms and charge. Do not generalise the simple net equation to every acid or every metal; the source specifies Mg, Zn and Fe with dilute hydrochloric/sulfuric acids.
A substance that neutralises an acid; it need not be soluble.
Use charges to make a neutral formula: Cu²⁺ and SO₄²⁻ give CuSO₄; Mg²⁺ and NO₃⁻ give Mg(NO₃)₂. CuO + H₂SO₄ → CuSO₄ + H₂O is balanced. CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂ has two chloride ions per calcium ion and one carbon dioxide molecule per carbonate formula unit.
Identify the acid first, then the positive ion and required ratio. Use brackets when more than one polyatomic ion is needed, keeping the nitrate or sulfate group intact. Confirm atom counts in the full equation. Observe carbon dioxide with the approved limewater test where appropriate, rather than treating all bubbles as hydrogen.
Formation of solid crystals from a solution under suitable concentration and cooling conditions.
Allow cooling and filter to remove excess insoluble oxide. The filtrate contains dissolved copper sulfate, so it passes through filter paper. Gently concentrate it in an evaporating basin using a water bath or electric heater, then allow cooling to form crystals. Separate the crystals and dry gently as directed. Evaporating completely to dryness is not the crystal-preparation endpoint.
Carry out the school-supervised task with the specified dilute reagents, eye protection and risk-assessed heat equipment. Record the actual sequence, observations and recovered dry sample. The handbook’s copper sulfate example uses warmed acid, then a water bath for evaporation. Avoid overheating the solution; hot apparatus can look cold. The teacher/technician determines appropriate crystal handling and disposal.
The scale used to express solution acidity or alkalinity.
Neutralisation between acid and alkali forms water: H⁺(aq) + OH⁻(aq) → H₂O(l). The positive and negative charges cancel, and atoms are conserved. Excess acid after mixing leaves the solution acidic; excess alkali leaves it alkaline. Neutralisation does not automatically mean that equal volumes of every pair of solutions will produce pH 7.
Use clean samples and small consistent indicator quantities or a correctly calibrated, rinsed pH probe. Match colours under suitable lighting and read the chart for the indicator actually used. During a supervised strong-acid/strong-alkali comparison record the measured pH against added volume. Indicator colour estimates and probe readings have different uncertainties; do not invent precision from a broad colour band.
A method measuring the volume of one solution needed to react with a measured amount of another.
Titre equals final burette reading minus initial reading. Read the appropriate meniscus at eye level, with the burette vertical and the filling funnel removed. A rough trial locates the endpoint; subsequent careful trials supply an accurate mean under the school’s stated concordance rule. Phenolphthalein is pink in alkali and becomes permanently colourless at the handbook’s acid-into-alkali endpoint. Universal indicator is unsuitable for a sharply defined accurate endpoint.
Perform the actual school-supervised task with approved dilute solutions and eye protection. Rinse apparatus as instructed to avoid unwanted dilution, fill the burette tip and use the pipette’s specified draining method. Never pipette by mouth. Retain initial/final readings, trials, selected mean and exclusion reason. Foundation students measure reacting volumes; concentration determination is the separate Higher extension.
The supplied concentration of a reference solution used in the calculation.
The acquired handbook PDF page 35 correctly states the equation and the 2:1 ratio, but prints a reversed half-mole statement afterwards. The correct inference is acid moles = half alkali moles. The balanced equation gives the correct amount relationship. The common-tier RP2 task still concerns accurate reacting volumes, while this numerical determination is Higher-only.
Use the measured accurate mean from the student’s actual supervised task, including the stated trial selection. Record known solution concentration and volume, convert cm³ to dm³, and label each chemical’s amount. If the question supplies different acid identity or coefficient ratio, change the calculation accordingly. A calculated concentration is not a direct burette reading.
The extent of ionisation of an acid in aqueous solution.
At a given concentration under comparable conditions, the stronger acid has lower pH because it produces greater hydrogen-ion concentration. Each decrease of one pH unit corresponds to a tenfold increase in H⁺ concentration. A decrease of two units means a hundredfold increase. Use whole-number pH comparisons in this source scope; do not treat pH as a linear concentration scale.
Compare labelled solutions of known identity and concentration using supplied data or teacher-approved dilute samples and a calibrated pH probe. Separate a strength comparison at fixed concentration from a concentration comparison using the same acid. The GCSE strong-acid wording is a model; do not silently turn it into advanced assumptions about every dissociation step of concentrated sulfuric acid.
A liquid or solution conducting charge through mobile ions.
Charge travels through the electrolyte by moving ions and through the external circuit by electrons. Positive ions are not positive electrons, and solid salt ions do not become mobile simply because a wire touches the solid. Electrolysis uses electrical energy to drive chemical change. The polarity names here apply to an externally powered electrolytic cell, not a blanket rule for every electrochemical cell.
Draw a labelled direct-current supply, two separated electrodes and the electrolyte. Mark supply polarity, electrode identity and arrows for each ion type. Actual school work uses approved low-voltage equipment and solutions, with electrodes held apart to prevent short circuits. Record products and observations at the named electrode instead of merely saying bubbles occurred somewhere.
A melted ionic compound whose ions can move.
Identify the positive metal ion and negative non-metal ion from the formula. The cathode product is the metal element, while a halogen anode product is diatomic, such as Br₂ or Cl₂. Do not write bromide as the elemental product. The formula’s ratio is determined by charge, and the product equation must preserve every atom.
Use diagrams, acquired observations or a teacher-approved demonstration rather than prescribing independent molten lead bromide work. The specification names anhydrous zinc chloride as a safer alternative, but it still needs an approved risk assessment and heated apparatus. Hot corrosive melts and toxic gas products cannot be treated like an ordinary room-temperature salt solution.
The material used in the molten mixture to lower the operating temperature of aluminium-oxide electrolysis.
Aluminium forms at the negative cathode. Oxygen associated with oxide discharge at the positive carbon anode reacts with carbon, forming carbon dioxide in the GCSE account. The carbon anode is consumed and needs continual replacement. It is therefore not an inert electrode like the one assumed in a simple product-prediction exercise.
Interpret a labelled industrial cell using supply polarity, electrolyte composition, metal collection and anode material. Use supplied electricity, temperature or electrode-replacement data to evaluate cost or resource demand. This is an industrial reference case, not a school task to build a hot extraction cell. Avoid adding unrequired plant details that obscure the two required explanations.
A water-containing ionic solution used as the conducting liquid.
Aqueous copper(II) chloride therefore gives copper at the cathode and chlorine at the anode. Sodium chloride solution gives hydrogen and chlorine under the approved handbook conditions. Copper(II) sulfate gives copper and oxygen, while sodium sulfate gives hydrogen and oxygen. State that these are the specified single-solute, inert-electrode predictions; concentration and electrode material can affect actual competing reactions outside this simple rule.
Identify the dissolved positive and negative ions, then include water-derived ions in the explanation. Apply cathode and anode rules separately. Use only school-approved dilute solutions and a low-voltage supply, with fixed electrode spacing and appropriate ventilation. Chlorine-containing observations must follow the technician-approved small-scale method; do not smell gases or generate large quantities.
A testable prediction with a stated reason.
The acquired handbook uses carbon rods, low-voltage direct current and approved CuCl₂/NaCl solutions. A brown/red deposit supports copper at the negative electrode; gas at the positive electrode bleaches damp blue litmus under the approved test. A bubble observation alone does not identify a gas. Label each observation with electrode sign and keep the products separate.
The teacher/technician approves concentrations, small quantities, voltage, duration, ventilation and disposal. Keep the carbon rods apart and fixed to avoid a short circuit. Use the prescribed contained small-scale gas test and never inhale chlorine. Record matched solution volumes, electrode spacing, supply setting and elapsed time. Retain actual results, labelled apparatus and a conclusion comparing them with the hypothesis.
An atom- and charge-balanced equation for one electrode reaction.
For 4OH⁻ → O₂ + 2H₂O + 4e⁻, oxygen count is four on each side and hydrogen count four on each side. Charge is −4 on both sides. For 2Cl⁻ → Cl₂ + 2e⁻, chlorine is diatomic and the electron total balances the two negative charges. A correct formula is needed before balancing; changing Cl₂ to Cl would misidentify the product.
Write the species and intended products first, then balance atoms, then charge with electrons. Check electrode sign and oxidation/reduction identity. Multiply a half equation only when matching overall electron transfer or the question’s required amount. Distinguish the basic metal/halide molten examples from water-derived aqueous examples and the industrial reacting carbon anode.
Describe metal reactions with oxygen forming metal oxides.
Identify oxidation and reduction by oxygen gain/loss without requiring the HT electron definition.
Recall room-temperature reactions of K, Na, Li, Ca, Mg, Zn, Fe and Cu with water/dilute acids.
Link metal reactivity to forming positive ions; include carbon and hydrogen in the series.
Predict displacement of a less reactive metal from its compound.
Deduce a reactivity order from experimental results while keeping comparison conditions explicit.
Explain native unreactive metals and extraction of less-reactive-than-carbon metals from oxides.
Identify oxygen gain/loss and evaluate a supplied extraction process without memorising industrial plant details.
Identify oxidation as electron loss and reduction as electron gain.
Write ionic displacement equations and match the electron numbers in the half equations.
Predict reactions of Mg, Zn and Fe with dilute hydrochloric and sulfuric acids.
Write balanced equations and distinguish acid–metal reactions from oxide/carbonate neutralisation.
Identify electron oxidation/reduction in acid–metal reactions.
Use metal and hydrogen-ion half equations to explain the balanced net reaction.
Predict acid reactions with hydroxides, oxides and carbonates.
Deduce salt identities/formulae using the acid and supplied positive/negative ion charges.
Describe preparing a named soluble salt from acid and an excess insoluble oxide/carbonate.
Explain heating, excess solid, filtration, concentration, cooling and drying as an actual supervised sequence.
Classify acidic, neutral and alkaline solutions with pH and use indicator/probe evidence.
Explain H⁺ and OH⁻ forming water during acid–alkali neutralisation.
Describe accurate strong-acid/strong-alkali titration with a suitable indicator.
Find titres and a defensible accurate mean; retain the common-tier practical separately from HT concentration calculations.
Calculate unknown titration concentration in mol/dm³ and convert to g/dm³.
Use the balanced acid–alkali ratio and independently check the practical handbook’s inconsistent Step 2 text.
Distinguish complete/partial ionisation from dilute/concentrated amount per volume.
Compare whole-number pH changes with tenfold H⁺ concentration changes and interpret relative acidity.
Explain electrolyte conductivity in molten/dissolved ionic substances.
Identify cathode/anode signs, ion movement and discharge products in electrolysis.
Predict products of a molten binary ionic compound with inert electrodes.
Distinguish the simple molten case from water-containing solutions and preserve practical suitability limits.
Explain using electrolysis for metals unsuitable for carbon reduction.
Explain cryolite lowering operating temperature and continual carbon-anode replacement.
Predict single-compound aqueous electrolysis products with inert electrodes using the GCSE rules.
Use metal reactivity relative to hydrogen and halide presence; distinguish prediction scope from all real conditions.
Develop a hypothesis and investigate aqueous electrolysis with inert electrodes.
Record electrode-specific observations and approved product tests; control comparisons and evaluate evidence.
Write/complete/balance electrolysis half equations throughout 4.4.3.
Identify cathode reduction and anode oxidation, including H⁺, OH⁻, halide and metal examples.

oxidation: Loss of electrons
reduction · ⁨还原⁩: Gain of electrons
titre: The volume delivered from a burette to reach the stated end-point
equivalence point · ⁨等当点⁩: The point of stoichiometric reaction completion
oxidation by oxygen: Oxidation identified as gain of oxygen
reduction by oxygen: Reduction identified as loss of oxygen
reactivity series: An ordering of metals by their tendency to react and form positive ions
positive ion: A particle with a net positive electrical charge
displacement reaction: A reaction in which a more reactive element replaces a less reactive one in a compound
comparison control: A comparison used to identify whether the intended treatment caused a change
native metal: A metal occurring naturally in its elemental form
metal extraction: Obtaining a metal from a naturally occurring source or compound
electron oxidation: Oxidation defined as loss of electrons
electron reduction: Reduction defined as gain of electrons
acid–metal reaction: A suitable metal reacting with acid to produce a salt and hydrogen
iron(II) salt: A salt containing Fe²⁺ ions
hydrogen ion: The H⁺ ion present in an aqueous acid in this GCSE model
net ionic reaction: A reaction representation showing participating ions while omitting unchanged spectators
base: A substance that neutralises an acid; it need not be soluble
alkali · ⁨碱⁩: A soluble base producing hydroxide ions in aqueous solution
crystallisation: Formation of solid crystals from a solution under suitable concentration and cooling conditions
excess insoluble solid: Undissolved reagent remaining after the acid has reacted under the preparation method
pH scale: The scale used to express solution acidity or alkalinity
neutralisation: Reaction of an acid with a base; H⁺ and OH⁻ form water for an acid–alkali reaction
titration: A method measuring the volume of one solution needed to react with a measured amount of another
concordant titres: Titre results agreeing within the stated practical tolerance
known concentration: The supplied concentration of a reference solution used in the calculation
titration concentration: A concentration inferred from reacting amounts and measured solution volumes
acid strength: The extent of ionisation of an acid in aqueous solution
partial ionisation: Ionisation of only some acid particles in the stated aqueous equilibrium
electrolyte: A liquid or solution conducting charge through mobile ions
cathode · ⁨阴极⁩: The negative electrode in the powered electrolysis model, where reduction occurs
molten electrolyte: A melted ionic compound whose ions can move
binary ionic compound: An ionic compound containing two elements
cryolite: The material used in the molten mixture to lower the operating temperature of aluminium-oxide electrolysis
anode · ⁨阳极⁩: The positive electrode in the powered electrolysis model, where oxidation occurs
aqueous electrolyte: A water-containing ionic solution used as the conducting liquid
inert electrode: An electrode not intended to react in the specified electrolysis
hypothesis: A testable prediction with a stated reason
electrode observation: A recorded change identified with its electrode sign and conditions
electrode half equation: An atom- and charge-balanced equation for one electrode reaction
discharge · ⁨放电⁩: Conversion of an ion at an electrode through electron transfer

4.5 · Energy changes

Transferring energy from reacting chemicals to the surroundings.
Distinguish the reacting chemicals from their surroundings. A thermometer in a reacting solution measures the temperature of that solution, which receives or supplies energy during the chemical change. A temperature rise supports an exothermic interpretation under the stated conditions; a fall supports endothermic behaviour. Heating a vessel externally can obscure this evidence. Not every process involving cooling is a chemical reaction: new substances must also be formed.
Evaluate a supplied warmer or cold-pack design against its purpose: suitable temperature, duration, risk of leakage, storage and single-use waste. Use information supplied for the particular product. Many self-heating cans exploit exothermic changes; some sports injury packs use endothermic changes. These examples explain energy transfers and do not recommend applying an untested chemical mixture to skin.
Final or peak temperature minus the stated initial temperature.
After each addition, replace the lid, stir gently and record the highest temperature. Repeat the whole investigation with fresh starting solutions and calculate means at each added volume. Plot mean temperature against total alkali volume added. Draw appropriate best-fit lines through the rising and falling regions and estimate their intersection. It estimates the peak between discrete readings; do not join every noisy point and call the highest measured point exact.
Use the school-approved risk assessment, eye protection and technician-prepared solutions. Sodium hydroxide presents a particular eye hazard. Keep initial acid volume, solution concentrations, starting temperatures, cup, stirring and reading procedure comparable. Record actual observations and an apparatus sketch in the laboratory; this lesson prepares and interprets that work, rather than certifying performance from a written answer. A lid and insulation reduce heat exchange but do not eliminate it.
The minimum energy that particles must have for a reaction to occur.
In an exothermic profile products lie below reactants, showing an overall energy decrease of the reacting chemicals and transfer to surroundings. In an endothermic profile products lie above reactants, showing energy taken from surroundings. Both profiles can have an activation barrier. The vertical difference between reactants and products is the overall change, distinct from the larger climb to the peak. Horizontal distance is not elapsed time or a measured reaction rate.
Draw axes first, mark clearly different reactant and product levels, then join them with a smooth curve rising above both. Label reactants, products and activation energy. Add a separate vertical arrow for the overall energy change. Compare the two diagrams using level differences, not the apparent length of the curve. Use relative energy units if numbers are supplied; they illustrate a model rather than requiring calorimetry.
Energy required to break a mole of the stated bonds in the supplied model.
Use the balanced equation and actual bond types. A coefficient multiplies all bonds in that molecule. For H₂ + Cl₂ → 2HCl, break one H–H and one Cl–Cl and form two H–Cl bonds. For 2H₂ + O₂ → 2H₂O, break two H–H bonds and one O=O bond, and form four O–H bonds. An O=O double bond uses its supplied double-bond value once; do not double a single-bond value.
Make separate broken and formed lists with counts, multiply each supplied value by its count, total each list, then subtract in the stated order. These values give an approximate model for the reaction quantities represented by the equation. Check units and sign and interpret the result in words. Do not infer this calculated value from a thermometer reading without a different model; AQA’s solution practical does not require that calculation.
A device whose chemical reactions produce electricity.
Use supplied readings or reactivity information to compare cells under specified conditions. Do not promise a universal voltage from a metal pair alone: electrolyte choice and conditions also matter. Connecting identical cells in series adds their potential differences in the ideal stated model; reversed orientation subtracts one cell’s contribution. Voltage measures potential difference, not the amount of chemical reactant remaining or the rate of all reactions.
For actual school work use approved metal strips and electrolytes, a high-resistance voltmeter and consistent exposed areas, separation, temperature and cleaning. Keep electrodes apart so they do not directly short together. Read the sign and magnitude with the stated meter connections. Do not attach a bench power supply to this generating-cell comparison or attempt to charge an unsuitable cell. Detailed commercial cell chemistry is not required by this section.
A cell whose intended use ends when its reactants are used up.
Choose a cell for its intended use. Compare required voltage, available operating time, mass, initial cost, replacement frequency, charging access and disposal information supplied. Repeated use can make a rechargeable option cheaper over time despite greater initial cost. A rarely used emergency device might place more weight on storage performance and immediate availability. These conclusions depend on stated data, not a universal rule that one cell type is always best.
Read the manufacturer’s stated type and use only the approved charger and school procedure for an actual recharge demonstration. Do not charge alkaline cells or construct an improvised charger. A written comparison can use fictional costs safely: state whether electricity, charger purchase and end-of-life disposal are included. Keep environmental and financial criteria distinct so a low purchase price is not automatically called lower waste.
A cell producing electricity from electrochemical reactions with externally supplied fuel and oxidant.
Compare both devices against a purpose. Supplied information may include mass, operating time, refuelling or charging time, infrastructure, fuel storage and cost. Water is the direct product of the hydrogen–oxygen reaction; hydrogen production and transport may have other environmental effects. A conclusion about total emissions needs data for those stages. Do not infer zero lifecycle carbon emissions merely from the absence of carbon in the local equation.
Use a teacher-approved educational fuel-cell kit or supplied observations; never improvise compressed hydrogen storage or a hydrogen–oxygen gas mixture. Record how fuel is supplied, where the electrical circuit connects and what product forms. A comparison task must use the given boundary: device-only, fuel production or a wider lifecycle. Detailed commercial engineering is not required, and the electron half equations are in the separate Higher lesson.
The electrode where hydrogen oxidation releases electrons in the stated generating cell.
Check atoms and charges separately. The hydrogen half equation has four hydrogen atoms on each side and zero net charge because +4 from H⁺ balances −4 from electrons. The oxygen half equation has two oxygen atoms and four hydrogen atoms on each side, again zero net charge. In this generating cell electrons leave the hydrogen anode through the external circuit and reach the oxygen cathode. The anode is negative and the cathode positive while the cell supplies electricity.
Begin with one consistent electrolyte model. Balance hydrogen with H⁺, oxygen with water and charge with electrons, then scale to equal electron numbers and cancel. Do not mix an acidic half equation with an alkaline one containing OH⁻. The electrode names identify oxidation at the anode and reduction at the cathode; signs must be interpreted for a generating fuel cell, rather than copied from an externally powered electrolytic cell.
Distinguish exothermic and endothermic reactions from energy transfers and temperature changes.
Recall combustion, oxidation, neutralisation and thermal decomposition examples; evaluate given uses.
Carry out an actual supervised investigation of a variable affecting temperature change in reacting solutions.
Record initial and highest temperatures, repeat measurements and interpret a temperature–volume graph.
Draw simple curved exothermic and endothermic profiles showing reactants, products and activation energy.
Explain activation energy as the minimum energy needed for reaction and classify profiles by relative levels.
Explain energy required to break bonds and released when bonds form.
Calculate overall changes using supplied bond energies and relate the balance to exothermic/endothermic behaviour.
Describe simple chemical cells and the effects of electrode/electrolyte choice on voltage.
Explain a battery as two or more cells connected in series and interpret supplied reactivity evidence.
Explain reactant exhaustion in non-rechargeable cells and reversal by external current in rechargeable cells.
Evaluate cells for particular uses with supplied cost, waste and performance information.
Explain hydrogen fuel cells using external hydrogen and oxygen supplies and the overall water-forming reaction.
Compare fuel cells with rechargeable cells/batteries using supplied evidence and system boundaries.
Write and balance the half equations for a hydrogen fuel cell.
Combine oxidation and reduction with equal electron counts to obtain the overall water-forming reaction.

exothermic · ⁨放热反应⁩: Transferring energy from reacting chemicals to the surroundings
endothermic · ⁨吸热反应⁩: Taking energy from the surroundings into reacting chemicals
temperature change: Final or peak temperature minus the stated initial temperature
best-fit intersection: Where fitted rising and falling trends meet to estimate a peak
activation energy: The minimum energy that particles must have for a reaction to occur
reaction profile: A diagram showing energy levels as a reaction progresses
bond energy: Energy required to break a mole of the stated bonds in the supplied model
overall energy change: Total bond-breaking energy minus total energy released by bond formation
chemical cell: A device whose chemical reactions produce electricity
battery: Two or more cells connected in series to provide a greater voltage
non-rechargeable cell: A cell whose intended use ends when its reactants are used up
rechargeable cell: A cell designed for its reaction to be reversed by an external current
fuel cell: A cell producing electricity from electrochemical reactions with externally supplied fuel and oxidant
lifecycle boundary: The stated production, use and disposal stages included in an evaluation
fuel-cell anode: The electrode where hydrogen oxidation releases electrons in the stated generating cell
fuel-cell cathode: The electrode where oxygen reduction accepts electrons in the stated generating cell

4.6 · Rate and extent of chemical change

The minimum energy that particles must have for a reaction to occur.
A product-time graph has a steeper gradient where rate is larger. A tangent estimates instantaneous rate; a secant gives average rate over an interval. The final plateau reflects the total collected product under the stated conditions.
For gas production, check apparatus for leaks, start timing consistently and record volume at regular intervals. Keep concentration, reactant amount and surface area controlled when changing temperature.
A state with equal forward and reverse reaction rates.
Beginning with mostly reactants, the forward rate may be high while the reverse rate is initially low. As products accumulate, reverse reaction becomes possible more often until the rates match under the stated model. A reaction-rate graph should show convergence to the same non-zero rate. A composition graph can level off at different reactant and product amounts. Identify which quantity is on the axis before claiming two curves must meet.
Use an original particle-count model or supplied data with fixed conditions and no matter lost. Count forward and reverse events over equal time windows. Their difference gives net change for a one-to-one A ⇌ B model. Do not build a sealed heated chemical apparatus to demonstrate the idea; safe school demonstration arrangements require approved containment and temperature control. A count model illustrates balance without proving a particular real chemical rate law.
Quantity of reactant used or product formed divided by the stated time interval.
Plot time on the horizontal axis and measured quantity on the vertical axis with labelled units and sensible scales. A steeper product-time curve means faster formation; a horizontal plateau means no further measured product forms. A tangent touches the local curve direction at the chosen point and its steepness indicates rate there. Drawing and interpreting tangents is both-tier; calculating their numerical gradients is in the Higher lesson. A straight chord between distant points gives an interval mean, not the local rate.
Collect readings at regular intervals with a consistent start. Check whether the graph is product formed, reactant used up or reactant remaining. Remaining reactant falls; its negative slope represents a positive rate of consumption when expressed as the amount used. Compare curves at the same time and under stated conditions. Do not assume a plateau proves every reactant is exhausted: a limiting reactant, incomplete collection or another stopped process needs consideration.
Rate at a stated time estimated by the gradient of a tangent.
First mark the requested time, draw a tangent following the smooth curve’s local direction, then select two readable points on that line. They need not be measured data points or lie on the curve elsewhere. A larger triangle reduces the relative effect of reading uncertainty. Tangent estimates can differ slightly between reasonable drawings. A chord linking two curve points measures an interval mean and must not be presented as the tangent gradient.
Show each coordinate pair, differences, division and final unit. For reactant remaining, the tangent gradient is negative; report its positive magnitude when asked for rate of reactant consumption. Convert time or amount units before dividing: milliseconds and seconds are not interchangeable, nor are millimoles and moles. This numerical tangent calculation and the molar-rate requirement are the embedded Higher clauses, while common-tier students still draw and interpret tangent steepness.
A condition whose change affects how quickly a reaction proceeds.
State exactly what changes and what is controlled. For equal masses of the same carbonate with acid in excess, smaller pieces can reach the same final carbon dioxide amount sooner. If increasing acid concentration also changes the limiting reactant amount, both speed and final amount can change; this is a different comparison. Higher gas pressure is relevant to reacting gases, not automatically a general explanation for every liquid reaction.
Choose one independent variable and measure a defensible rate response: initial curve steepness, quantity formed in a fixed time, or time to a defined comparable endpoint. Keep other factors controlled, including temperature, total liquid volume, solid mass and exposed surface where appropriate. Repeat measurements and retain observations. A catalyst must be suitable for that reaction; a biological enzyme is not a universal catalyst for unrelated chemical changes.
Monitoring gas product volume over time to investigate reaction rate.
Fit the collector and check for unobstructed movement and an open gas path into the collector. Add magnesium, replace the bung promptly and start timing consistently. Record gas volume at suitable intervals, for example every 10 s, until it becomes constant. Plot both volume–time curves on the same axes and compare steepness at comparable times. Identical final volumes require equal limiting magnesium and enough acid; real collection losses or sample differences can change recorded plateaux.
Perform the approved actual school experiment with eye protection, a stable clamped collector and no ignition sources near hydrogen. Use matched ribbon mass/length and consistent oxide removal, acid volume, initial temperature and apparatus. Never seal gas-producing apparatus with no outlet. Repeat trials where feasible and compare methods. A syringe avoids some water-collection problems but can stick or leak; a water collector needs consistent reading conditions and gas not excessively soluble in water.
Cloudiness caused by suspended particles affecting how clearly a target is seen.
The concentration labels refer to the 50 cm³ pre-acid dilution; adding acid further dilutes the mixture consistently. Maintain total volumes, flask, viewing depth, lighting, temperature and endpoint judgement. Repeat and calculate mean endpoint time. For a comparable fixed turbidity threshold, 1/mean time is a relative rate proxy: a shorter time means a larger proxy. It does not measure a known mass of sulfur per second or the exact completion time of the entire reaction.
Carry out only the school-approved method with ventilation and eye protection. Sulfur dioxide is released and can aggravate breathing difficulties; do not lean over the vessel or deliberately inhale fumes. The teacher selects controls and an appropriate alternative observation arrangement for affected students. Record raw times and reasons for any justified anomalous-result exclusion; do not discard values merely because they spoil a trend. Compare repeats within a method before claiming reproducibility across groups.
How often reacting particles collide under the stated conditions.
Keep frequency distinct from collision energy. At unchanged temperature, raising concentration alone does not mean particles move faster or each has more energy. In a simple stated model with one collision partner held fixed, doubling the other partner’s particle concentration can double encounter opportunities. It is not a universal measured rate law for every mechanism or for simultaneous changes to both reactants. Different reactions require suitable evidence for exact proportional relationships.
Compare same-volume particle drawings using equal particle sizes and the same stated temperature. Count only the labelled reacting species, not solvent particles as if they were all reactants. For gas compression specify that temperature is controlled; actual compression can otherwise change temperature too. School experiments use approved dilute solutions and supplied gas data rather than student-built pressure vessels. Translate a rate claim into collision frequency and sufficient-energy reasoning.
A collision carrying enough energy to meet the reaction’s activation requirement.
A cube model makes the size comparison explicit. A cube side L has area 6L² and volume L³. Cutting it into eight half-side cubes preserves total volume but doubles total area. This model assumes newly exposed faces are accessible and not clumped or coated. Real powders need appropriate containment; aggregate formation can reduce accessible area. Do not claim every ten-degree temperature rise always doubles rate: exact effects depend on reaction and conditions.
Investigate one factor at a time with teacher-approved dilute reactants and moderate controlled temperatures. Use a water bath where appropriate and measure the reacting mixture’s temperature, not only the bath label. Compare equal solid masses with the same composition and control acid concentration and volume. Fine powders can create exposure hazards; use the approved particle sizes rather than grinding unknown materials independently. Explain the measured trend through the relevant collision mechanism.
A substance increasing reaction rate by a lower-activation pathway without being used up overall.
At the same temperature more colliding particles can meet the smaller activation requirement. This does not mean the catalyst adds energy to every particle or increases temperature by definition. Evidence for catalytic behaviour combines rate increase and regeneration/no net consumption. An unchanged small sample alone does not prove catalysis unless the reaction-rate effect is shown. The overall equation omits a regenerated catalyst from its stoichiometric reactant and product lists.
Compare teacher-approved catalysed and uncatalysed trials with matching reactant quantities, concentrations, temperature and collection apparatus. Record time-dependent product measurements rather than only final volume. Use the specified catalysts in the course where required; do not memorise unrelated industrial names. Actual hydrogen peroxide catalyst demonstrations need the school’s selected dilute solutions and protection because rapid gas evolution and warming can occur.
A reaction whose products can react to reform the original reactants.
The symbol does not mean that all substances must be present in equal quantities or that the reaction continually switches completely from one side to the other. A reversible reaction may reach equilibrium when substances are contained under suitable fixed conditions, but reversibility alone is not proof that an open apparatus has reached equilibrium. Melting/freezing is a reversible physical change; it does not by itself demonstrate this chemical product-to-reactant relationship.
Use supplied observations or a teacher demonstration under the approved risk assessment. Ammonia and hydrogen chloride gases are irritant; students must not generate or smell them independently. Record where gas forms and where the white solid reappears, distinguish observation from chemical explanation and retain the condition labels. Do not improvise a sealed heated vessel. Write a balanced equation with suitable state symbols only for the stated conditions.
The blue copper sulfate material containing water of crystallisation in this example.
Products with greater energy in the forward endothermic reaction become the higher-energy starting materials for the reverse exothermic reaction. Reversing the equation swaps the levels. The equal transferred amounts do not mean equal activation energies for the two directions: the climb to the same peak starts from different levels. That distinction can be read qualitatively without introducing an advanced thermodynamic calculation.
Use teacher-approved demonstration samples or supplied before/after evidence. Copper salts need controlled handling and waste collection; adding water to hot material or heating unknown compounds is inappropriate. Record colour alongside the identified substances and temperature evidence. Water loss from a hydrate is a specific chemical change, not a rule that every blue substance becomes white on heating or that every colour change is reversible.
A closed-system state with equal continuing forward and reverse reaction rates.
An equilibrium system responds to a disturbance in a direction that counteracts it.
Begin with the system already at equilibrium, identify one imposed change and select the appropriate response rule. Increased reactant concentration favours its consumption; removing product favours forming more. Warming favours the endothermic direction. Compression of a gaseous equilibrium favours the side with fewer gas molecules. These are composition responses, not the same as saying every faster reaction makes a larger final yield.
Use supplied equations, forward energy direction, states and before/after data. Mark a shift arrow as net adjustment toward a new equilibrium, not cessation of the other reaction. Keep other conditions specified. A suitable catalyst increases both rates and reduces time to equilibrium without shifting the final equilibrium composition at fixed conditions. No equilibrium-constant expression or numerical constant calculation is required by this GCSE section.
An imposed change to a reactant or product concentration at equilibrium.
Separate the immediate imposed change from the response. If product is removed, its amount first falls; a subsequent net forward change can raise it from that lowered value without necessarily restoring the original level. At the new equilibrium both directions again occur at equal rates. A plot may show a sudden jump in the changed species followed by slower adjustment; the other species need not have the same immediate jump.
Use original fictional particle or concentration data with a stated fixed temperature and appropriate mixture model. Identify the species directly changed, then predict which direction uses it or replaces it. Real solution mixing can change volume as well as amount, so concentration comparisons need that information. The simple one-to-one particle model is an explanation aid; it is not a numerical equilibrium-constant method.
The reaction direction taking energy from the surroundings.
Both forward and reverse reactions can become faster on warming, yet their relative rate balance changes, causing a net shift before a new equilibrium is reached. For an exothermic forward industrial reaction, a low temperature may favour yield while slowing production; an appropriate temperature choice can be a compromise. A catalyst can help rate without turning that exothermic forward direction into an endothermic one or increasing its equilibrium yield at fixed conditions.
Read a supplied equation and mark forward endothermic/exothermic before writing a shift conclusion. Use data collected after equilibrium at each stated temperature, with other relevant conditions specified. A temporary product increase during heating is not necessarily the final equilibrium composition. This lesson predicts direction and evaluates supplied evidence; no numerical equilibrium constant or thermodynamic derivation is required.
The sum of balanced equation coefficients for gaseous species on one side.
If the gas-molecule counts are equal on both sides, changing pressure by compression gives no pressure-driven composition shift in this model. H₂(g) + I₂(g) ⇌ 2HI(g) has two on each side. If solids or liquids occur in a supplied equation, do not count their coefficients as gas molecules. Compression can also affect rates by increasing gas particle concentration; a rate effect alone is not proof of an equilibrium shift.
Write state symbols and a gas-count total under each side. Specify compression or expansion of the reacting mixture at constant temperature: merely adding an unrelated gas does not justify blindly applying the same rule without a volume/pressure description. Evaluate a supplied industrial choice using yield, rate, equipment cost and operating constraints. Do not ask students to construct high-pressure reaction vessels.
Calculate mean rate from mass of reactant used or volume of product formed per time.
Draw and interpret quantity–time graphs and tangent steepness without requiring HT numerical gradients.
Calculate the gradient of a tangent as reaction rate at a stated time.
Use quantity in moles and express rates in mol/s.
Recall effects of concentration, gas pressure, solid surface area, temperature and catalysts on rate.
Separate a faster reaction from an increase in final product amount.
Develop a hypothesis and carry out the gas-volume method of RP5 under school supervision.
Control magnesium samples/acid volumes, collect time-series readings and evaluate leaks/start delay.
Develop a concentration hypothesis and perform RP5 colour/turbidity investigation with supervised observations.
Use controlled depth/lighting, repeats and a qualified reciprocal-time rate proxy.
Explain concentration/pressure effects through collision frequency and sufficient collision energy.
Use proportionality only under the explicitly stated particle comparison assumptions.
Explain temperature effects using collision frequency and more energetic collisions.
Explain smaller-solid effects using exposed surface-area-to-volume ratio.
Identify catalysts from increased rate, no overall consumption and absence from the overall equation.
Explain different lower-activation pathways, reaction specificity and enzymes as biological catalysts.
Interpret the reversible arrow and explain how changing conditions changes reaction direction.
Use the specified ammonium chloride heating/cooling example without treating melting as reaction.
Explain opposite exothermic/endothermic directions with equal transferred amounts for a reversed change.
Interpret blue hydrated/white anhydrous copper sulfate and water as the specified example.
Describe equilibrium in an apparatus preventing loss of reactants and products.
Distinguish equal forward/reverse rates from equal amounts or stopped reactions.
Use Le Chatelier’s Principle to predict a qualitative equilibrium response to a stated change.
Distinguish restoring a rate balance from completely undoing the imposed change.
Predict equilibrium responses to specified reactant/product concentration changes.
Interpret supplied composition evidence while separating immediate disturbance from later adjustment.
Predict temperature effects using the stated forward endothermic/exothermic direction.
Interpret given equilibrium composition data without confusing rate increase with equilibrium yield.
Predict compression/expansion effects using the number of gaseous molecules on each side.
Interpret states and equal-gas-count reactions without counting solids as gas.

activation energy: The minimum energy that particles must have for a reaction to occur
rate · ⁨率⁩: Change in a measured quantity per unit time
equilibrium · ⁨平衡⁩: A state with equal forward and reverse reaction rates
reversible reaction: A reaction whose products can react to reform the original reactants
mean reaction rate: Quantity of reactant used or product formed divided by the stated time interval
tangent · ⁨相切⁩: A straight line matching a curve’s local direction at a chosen point
instantaneous rate: Rate at a stated time estimated by the gradient of a tangent
molar rate: Amount of reactant used or product formed in moles per unit time
rate factor: A condition whose change affects how quickly a reaction proceeds
surface area: The exposed area of a solid available for contact with reacting particles
gas-volume method: Monitoring gas product volume over time to investigate reaction rate
start delay: A delay between reaction beginning and measurement or collection starting
turbidity: Cloudiness caused by suspended particles affecting how clearly a target is seen
reciprocal-time proxy: One divided by time to a comparable fixed endpoint, used as a relative rate measure
collision frequency: How often reacting particles collide under the stated conditions
energetic collision: A collision carrying enough energy to meet the reaction’s activation requirement
surface-area-to-volume ratio: Exposed surface area divided by the solid’s volume
catalyst · ⁨催化剂⁩: A substance increasing reaction rate by a lower-activation pathway without being used up overall
catalysed pathway: An alternative reaction route with a lower activation barrier
reversible arrow: Opposed half arrows indicating a reaction can proceed in both directions
hydrated copper sulfate: The blue copper sulfate material containing water of crystallisation in this example
anhydrous copper sulfate: The white copper sulfate material without water of crystallisation in this example
dynamic equilibrium: A closed-system state with equal continuing forward and reverse reaction rates
closed system: A system preventing escape or entry of reactants and products during the stated comparison
Le Chatelier’s Principle: An equilibrium system responds to a disturbance in a direction that counteracts it
equilibrium shift: A net composition adjustment toward a new equilibrium after conditions change
concentration disturbance: An imposed change to a reactant or product concentration at equilibrium
re-equilibration: Adjustment until forward and reverse reaction rates become equal again
endothermic direction: The reaction direction taking energy from the surroundings
exothermic direction: The reaction direction transferring energy to the surroundings
gas-coefficient count: The sum of balanced equation coefficients for gaseous species on one side
compression: Reducing the volume of a stated gas mixture to raise its pressure at fixed temperature

4.7 · Organic chemistry

An atom group determining characteristic reactions.
Distinguish addition, substitution, oxidation and polymerization by tracing bonds before and after reaction. Conditions and reagents belong to the reaction arrow; they are not interchangeable labels.
Draw displayed formulae with correct carbon and hydrogen valencies. For the specified alkene additions, keep the carbon skeleton and track the atoms added across C=C. Compare the supplied structure before and after reaction. GCSE 8462 does not require an advanced multistep synthesis route in this focus.
A finite mixture of many compounds, mainly hydrocarbons, formed from ancient biomass.
A mixture contains different substances together without a fixed bonding ratio between those substances. Each individual compound still has its own definite formula. C₃H₈ is a hydrocarbon; C₂H₅OH contains oxygen and therefore is not a hydrocarbon, even though it contains carbon and hydrogen. Carbon dioxide is not a hydrocarbon either. Finite describes a limited resource, not a prediction that every deposit will disappear on one exact date.
Read every element symbol before classifying a formula. Then distinguish statements about an individual molecule from statements about the whole mixture. Use supplied refinery composition data to discuss which components are present and their proportions. A classroom model can use labelled molecular cards; it does not require obtaining or heating crude oil. Keep origin, composition and use as separate parts of an explanation.
A saturated hydrocarbon in the homologous series with general formula CnH2n+2.
For a straight chain, each end carbon has three hydrogen atoms, while an internal carbon has two. Methane has no carbon neighbour and so has four hydrogens. Count the bond lines at each carbon to check a displayed structure rather than adding hydrogen randomly. The general formula identifies the expected molecular composition of this series; molecular composition alone need not establish a unique arrangement of atoms.
Construct the first four molecules using a model kit or the original displayed diagrams. Count carbons, total hydrogens and bonds at each atom. Translate between displayed and molecular formulae, then test a supplied formula against H=2n+2. A question may supply a longer-chain formula without requiring its specific name. Keep the required names limited to the four in the acquired section.
A separated mixture of hydrocarbons with similar carbon numbers and boiling ranges.
Evaporation and condensation are physical changes: molecules are separated without breaking their carbon chains. A sufficiently volatile fraction can remain gaseous at the top; heavy material can remain near the base. Fuels include petrol, diesel oil, kerosene, heavy fuel oil and liquefied petroleum gases. Fractions also provide feedstock for the petrochemical industry, producing materials such as solvents, lubricants, polymers and detergents. Carbon atoms can bond to each other in families of related structures; this helps explain the wide variety of natural and synthetic carbon compounds.
Trace a labelled vapour through the temperature-gradient diagram and explain where it condenses using its supplied boiling range. Compare a fraction’s intended fuel use with a feedstock use. Do not replace an explanation with a memorised order of names alone; names of other specific fractions are not required. School demonstrations must use an approved substitute apparatus and risk assessment, not a sealed heated crude-oil container.
Resistance of a fluid to flow.
Keep the three comparisons separate. Boiling point concerns liquid becoming gas, viscosity concerns flow, and flammability concerns ignition and burning. Stronger intermolecular attractions in larger comparable hydrocarbon molecules help explain higher boiling points. Boiling does not normally break the covalent bonds within the molecules. The acquired section limits recalled property trends to these three; density or a precise energy-per-molecule trend is not required here.
Use supplied property data to select a fuel and justify it against a stated purpose. A ranking is meaningful only if temperature and test conditions are comparable. A teacher-supervised investigation can compare flow times or supplied ignition observations using approved small samples and controlled equipment. Students must not ignite unknown fuels independently, and a flame comparison alone cannot establish every property.
Burning a hydrocarbon in sufficient oxygen to form carbon dioxide and water.
One carbon atom requires one carbon dioxide molecule, while each pair of hydrogen atoms forms one water molecule. Oxygen atoms on the product side come from both products, so count both before finding O₂. If an intermediate oxygen coefficient is a half number, multiply every coefficient by two for smallest whole-number coefficients. Never alter the given hydrocarbon’s subscripts to force balance.
Write correct formulae first, balance carbon and hydrogen, then oxygen, and check all three element totals. State the oxygen condition in the explanation. Supplied combustion observations can support energy release; they do not alone prove every carbon atom became carbon dioxide. Incomplete combustion products and atmospheric effects belong to their other course sections. Use teacher-approved demonstrations or written evidence rather than independent fuel-burning trials.
Breaking larger hydrocarbon molecules into smaller, more useful molecules.
High demand for small-molecule fuels makes some cracking products useful as fuels. Alkenes can be used to make polymers and other chemicals. A cracking equation is an atom-conservation example; different products can form depending on conditions. A formula supplied in the equation must be preserved. The full named alkene structures and addition reactions are separate Chemistry-only 4.7.2 requirements, not evidence that this introductory lesson completes them.
Count carbon and hydrogen on each side of the supplied equation, solve for a missing formula or coefficient, and check both elements. Contrast cracking with distillation: cracking forms new substances, while distillation separates existing ones. Bromine-water observations need approved teacher-controlled microscale equipment and ventilation; students must not improvise a heated cracking apparatus or expose themselves to bromine. Written interpretation prepares practical reasoning without certifying laboratory work.
An unsaturated hydrocarbon containing a carbon–carbon double bond.
A displayed formula shows all atoms and all bond lines. Count a double bond as two of carbon’s four bonds. In the terminal-double-bond structures shown, the first carbon has two hydrogens and the next has one, except ethene where each has two. A longer chain finishes with CH₃. Butene and pentene can have different atom arrangements; the diagrams show one stated straight-chain arrangement, rather than claiming their molecular formula uniquely fixes a structure.
Build the four illustrated molecules using a model kit, then translate each into a displayed drawing and molecular formula. Check hydrogen valency one and carbon valency four. Explain what a flat drawing reveals about connectivity and what it omits about the real three-dimensional molecule. Recognise a supplied longer member from its double bond and formula without requiring a specific name beyond the four stated members.
Addition of hydrogen across a carbon–carbon multiple bond to form a more saturated product.
An addition reaction joins atoms to the molecule across the multiple bond without removing the existing carbon skeleton. Check the new bond counts: each carbon still makes four bonds. Separately, alkenes burn in oxygen like other hydrocarbons. With sufficient oxygen, complete combustion gives carbon dioxide and water. In air they tend to burn with smoky flames because incomplete combustion produces carbon particles; smoke does not establish complete combustion.
Draw the given alkene first, replace C=C by C–C and add one H to each affected carbon. Retain every other atom. Use the individual displayed product panels for all four named starting members. For combustion, balance carbon, hydrogen and then oxygen. Analyse approved demonstration evidence rather than conducting a pressurised hydrogen/catalyst experiment: combustible gases, heated catalysts and oxygen must remain under trained control.
Addition of one halogen atom to each carbon of a carbon–carbon double bond.
For ethene plus bromine, CH₂=CH₂ becomes BrCH₂–CH₂Br. Each carbon retains its two hydrogen atoms and gains one bromine, giving four bond contributions. Bromine water changes orange to colourless in the specified reaction; an alkane control does not give the same rapid addition under these conditions. A carbon skeleton and hydrogen count are retained while two halogen atoms are added. The displayed panels distinguish chlorine, bromine and iodine products for all four named starting alkenes.
Draw the original molecule, mark the two double-bonded carbons, change their link to single and attach one halogen to each. Check every carbon and every hydrogen rather than replacing an H with Br. Record actual colour against a blank/control in teacher-approved microscale demonstrations. Halogens need approved containment, eye protection and ventilation; students must not generate chlorine, handle elemental bromine or improvise iodine reaction conditions independently.
Addition of water across a multiple bond in the stated alkene reaction.
The product is an alcohol, not a hydrocarbon, because it now contains oxygen. Count the H inside OH as well as those bonded directly to carbon. Ethene’s two carbons are equivalent in this simple addition. Unsymmetrical starting alkenes can give different OH positions: the longer terminal-chain panels show the usual secondary-alcohol arrangement with OH on the second carbon. These drawings identify stated products without adding a required advanced mechanism or demanding positional product names beyond the course scope.
Draw all bonds including C–O and O–H, and check carbon four, oxygen two and hydrogen one. Compare the product with the input molecule and steam rather than adding OH alone and losing the extra H. The paired first-four displays show before/after connectivity. These industrial conditions are reference information, not a school pressurised synthesis task. Actual school work can use models and supplied process evidence under supervision.
An organic compound containing an –OH functional group attached to its carbon framework.
The first-four alcohols are used as fuels, solvents and starting materials for making other chemicals. Methanol is an industrial fuel/feedstock and solvent; ethanol is used as a fuel and solvent in suitable formulations; propanol and butanol serve as solvents and chemical feedstock. The selected use must be tied to actual product information, concentration and purpose. None of these industrial uses implies a sample is suitable to drink or apply to skin.
Compare molecular, condensed and displayed formulae. Locate OH, count the carbon chain and identify the given member. Distinguish a single alcohol molecule from an aqueous solution containing many water molecules as well. Use teacher-provided labelled structures or approved product evidence; do not identify a liquid by taste or smell. Specific names beyond the four stated members are not required for the acquired section.
Reaction of the stated alcohol with an oxidising agent under the given conditions.
For the displayed primary examples, methanol can give methanoic acid, ethanol ethanoic acid, propanol propanoic acid and butanol butanoic acid. The OH carbon’s position matters; a different alcohol arrangement cannot automatically be assigned the same oxidation product. In an approved test, acidified potassium dichromate(VI) can change orange to green as it oxidises the primary alcohol. The acquired GCSE requirement asks for observations/descriptions, not aldehyde mechanisms or balanced equations for these non-combustion reactions.
Record reagent, conditions and observation in separate columns. Describe solution formation with water as mixing/dissolving, not proof that water has made a new alcohol. Any sodium or oxidising-agent demonstration requires the school’s risk assessment, trained handling, eye protection and approved small quantities; chromium(VI) reagents need controlled waste. Students should analyse reference observations if the school does not approve the demonstration. Do not add sodium to an aqueous alcohol mixture to infer an alcohol-only reaction.
Anaerobic conversion of sugars by yeast to an aqueous ethanol solution and carbon dioxide.
Maintain the school-selected temperature and sugar concentration, use the prepared yeast and record actual gas or mass observations over time. An airlock can let gas leave without freely admitting outside air; it is not a licence to tightly seal a gas-producing bottle. The rate can change as substrate runs low or ethanol inhibits yeast. Gas bubbling is qualitative unless quantity is measured, and a silent vessel is not by itself proof that no ethanol is present.
For an actual supervised activity, compare matched labelled mixtures with a yeast-free control and the same volumes, sugar concentration and temperature. Retain raw observations and repeat comparisons. Analyse differences cautiously because gas leakage, dissolved carbon dioxide and contamination affect measurements. Any later separation of ethanol requires approved equipment and no flame near flammable vapour. Products of school fermentation must not be consumed.
An organic acid containing the –COOH functional group.
These first-four acids dissolve in water to make acidic solutions. A suitable indicator or measured pH demonstrates acidity under the stated conditions. They react with carbonates to form a salt, carbon dioxide and water. For example, ethanoic acid with sodium carbonate forms sodium ethanoate plus the gas and water. Carbon dioxide can turn limewater milky in an approved product test. Reaction with alcohols forms an ester and water, developed in the separate ester lesson.
Compare a supplied acid structure with the named member by counting all carbons, including COOH. For an approved supervised carbonate comparison, use teacher-selected dilute acids, comparable volumes and carbonate quantities, and record effervescence or gas measurements. Keep a path for gas to escape or reach the approved collector. Indicator colour identifies acidity, not the precise acid; different concentration or temperature can affect reaction rate.
An organic compound containing the C(=O)–O–C connection formed here from an acid and alcohol.
The illustrated product contains a C(=O)–O–C connection. The acid-derived carbon framework and alcohol-derived carbon framework both remain. It is different from simply mixing an acid with water, from sodium carbonate reaction and from alkene addition. Formation need not be complete in one trial. A characteristic smell may be supplied as reference evidence but is not a unique identification test or a reason to inhale concentrated vapour.
Use the word equation and the displayed model to trace carbon and oxygen atoms. In an approved teacher-led demonstration, record the starting reagents, warming/catalyst conditions and evidence without treating a fruity description alone as conclusive. A comparison may use supplied spectra or labelled products if no practical is approved. Do not improvise a flame, sealed heated vessel, concentrated acid addition or consumption of the product. Actual school evidence remains separate from written practice.
The extent of ionisation of an acid in aqueous solution.
Concentration describes the amount of acid in a volume; strength describes its extent of ionisation. These are independent labels. Comparing a dilute strong acid with a concentrated weak acid does not establish the simple same-concentration pH ordering. For whole-number pH differences, a decrease of one pH unit means ten times the hydrogen-ion concentration. No equilibrium constant, pKa calculation or advanced weak-acid approximation is required here.
Use teacher-provided concentrations and pH data or an approved supervised pH measurement. Keep temperature and measurement calibration comparable. State what is measured and what explanation depends on equal concentration. A particle model should conserve total acid-derived units and electrical charge: each ionised unit supplies one positive H⁺ and one negative carboxylate in the illustrated monoprotic example. Particle counts are a model, not a measured degree of ionisation for all samples.
A small molecule that can join with others to form a polymer.
Draw the two-carbon backbone horizontally, keep both hydrogens on each carbon and show a single bond extending out of each side of the brackets. Put n outside the brackets to represent many repeats. A continuation bond joins the neighbouring unit; it is not an extra hydrogen or a separate monomer still containing C=C. Carbon has four bonds when the continuation bonds are counted. The bracketed unit represents the repeating interior of the chain, not its complete end-group structure.
Use a molecular kit or atom-labelled cards to build several ethene molecules. Replace each C=C by a single link, then join their carbon backbones without removing any atoms. Translate that model into a short chain segment and a bracketed repeating unit. Reverse the reasoning for a supplied unit: identify the adjacent backbone carbons that came from the double bond, restore C=C and remove the continuation bonds when drawing the monomer.
The smallest stated structural pattern repeated along the polymer chain.
Trace the two C=C carbons first and label their attached atoms or groups before changing any bond. In the repeat, the CH₂ carbon bonds to two hydrogens and two backbone neighbours. The CH carbon bonds to one hydrogen, the CH₃ group and two backbone neighbours. Both have four bonds. A supplied CH₂=CHCl example similarly gives –CH₂–CHCl–: chlorine stays attached. This supplied structure is for transferring the model, not an extra list of commercial names to memorise.
Draw a given alkene in a fixed orientation, circle the double-bond carbons and copy their side groups into the repeating unit. Replace the double bond with a single one, add continuation bonds and brackets, then check every atom and valency. For a reversed task, remove the continuation bonds and restore C=C between the corresponding backbone carbons. Equivalent reversed drawings describe the same connectivity; placement on the page is not a different polymer identity.
Joining multifunctional monomers while forming small-molecule by-products in the stated reaction.
Read the ideal repeating interior as –O–CH₂–CH₂–O–C(=O)–(CH₂)₄–C(=O)–. The CH₂–CH₂ segment came from the diol, and the four-CH₂ segment and carbonyl carbons came from the diacid. At each joining event, H from an alcohol OH and OH from an acid COOH make water. The remaining oxygen links the carbonyl carbon to the diol segment. Unlike addition polymerisation, the repeat therefore does not retain every atom of the original monomer pair.
Build a paper or model-kit chain from two-ended diol and diacid cards. Join only compatible ends, mark each ester link and place one water card beside each actual joining event. A single OH/COOH pair cannot extend indefinitely at both ends. Compare the repeat with both starting molecules and explain which functional groups enabled growth. This modelling task is not permission to perform an unapproved polymer synthesis in the laboratory.
The nitrogen-containing –NH₂ functional group in the stated amino-acid structure.
Each joining event loses OH from one carboxylic-acid group and H from one amino group as H₂O. A chain can extend because functional groups remain at its ends. Different amino acids can combine in the same chain to make proteins; their order and folding affect the resulting protein. Glycine-only repetition is a simple model, not a claim that every natural protein contains one repeating amino-acid type. Detailed zwitterion, stereochemistry and protein-structure classifications are not required here.
Use labelled amino-acid cards with distinguishable amino and acid ends. Form a short open chain, mark each peptide link, retain the carbonyl double bonds and account for each water molecule. Then build a chain using several different amino-acid labels to represent sequence variation. Keep chemistry evidence separate from a dietary claim: identifying a peptide bond does not establish the biological role, nutritional value or safety of an unknown sample.
A monomer building block of a nucleic-acid chain, containing a sugar, phosphate and a base.
The chains provide a repeated structural framework while the sequence of nucleotide types varies. A ladder model can distinguish the two continuous chains from paired positions across them. In the familiar double-stranded DNA model, A pairs with T and C with G. Twisting that ladder models the double helix. Colour-coded blocks are a representation: they are not actual nucleotide shapes, accurate atom arrangements or a diagram of separate bases joined without a backbone.
Build an actual classroom model with four labelled nucleotide-card types. Arrange a short sequence along one chain and the complementary sequence along the other, then twist or compare with a supplied three-dimensional model. State what the model shows about repeated units and variable sequence, and what it omits about dimensions and chemical bonds. No human DNA collection or biological sampling is necessary for this representation task.
A polymer produced in a biological system.
Proteins can use different amino acids in the same chain, with sequence and folding contributing to their functions. Starch and cellulose use glucose but differ in how their building blocks are joined and arranged, producing different properties. Starch stores carbohydrate in plants; cellulose gives strength to plant cell walls. The required Chemistry comparison is monomer identity and polymer recognition, not memorising advanced alpha/beta linkage names or complete structural formulae of carbohydrates.
Sort labelled structure or identity cards into monomer/polymer pairs, then justify each match aloud using the building-block evidence. Place starch and cellulose in separate polymer groups even though their glucose cards match. If using a teacher-approved starch observation with iodine, record the actual sample and observed colour; a result for starch does not identify every glucose-containing polymer. Keep practical observations, structural inferences and biological-use statements in separate columns.
Describe the ancient-biomass origin and finite nature of crude oil found in rocks.
Identify hydrocarbons as compounds containing hydrogen and carbon only.
Use CnH2n+2 for the alkane homologous series and name methane, ethane, propane and butane.
Recognise alkane molecular and displayed formulae without requiring other specific alkane names.
Explain how crude oil fractions separate using evaporation, condensation and a temperature gradient.
Explain fuel and petrochemical feedstock uses without treating a fraction as one pure compound.
Recall increasing boiling point and viscosity and decreasing flammability with increasing molecular size.
Apply those three trends to supplied fuel-use evidence without requiring extra property trends.
Explain oxidation of carbon and hydrogen and formation of carbon dioxide and water in complete combustion.
Balance complete-combustion equations for supplied hydrocarbon formulae.
Describe general catalytic and steam-cracking conditions and alkane/alkene products.
Balance supplied cracking equations, recall the bromine-water test and explain fuel/feedstock uses.
Use CnH2n and recognise ethene, propene, butene and pentene in molecular and displayed forms.
Explain unsaturation using C=C and two fewer hydrogens than the corresponding alkane.
Describe hydrogen addition using a heated nickel catalyst and draw the saturated product.
Explain smoky alkene flames in air and balance complete combustion when sufficient oxygen is supplied.
Describe room-temperature halogen addition and draw all atoms in the products for the first four alkenes.
Explain orange bromine-water decolourisation in the stated alkane/alkene comparison.
Describe hydration with steam, high temperature/pressure and a phosphoric-acid catalyst.
Draw displayed hydration products while checking which carbon receives H and which receives OH.
Recognise –OH in methanol, ethanol, propanol and butanol from names and formulae.
Recall alcohol uses as fuels, solvents and chemical feedstock, linking function to stated properties.
Describe sodium reaction, combustion, mixing with water and oxidation for the taught primary alcohols.
Balance alcohol combustion, without requiring equations for the other alcohol reactions.
Describe fermentation of sugar solutions with yeast under warm anaerobic conditions.
Distinguish aqueous ethanol, gas evidence and supervised fermentation from the industrial steam route.
Recognise –COOH in methanoic, ethanoic, propanoic and butanoic acids.
Describe dissolving in water and carbonate reactions without requiring balanced acid-reaction equations.
Describe ester formation from a carboxylic acid and alcohol and recognise water as the other product.
Name ethyl ethanoate from ethanol/ethanoic acid without requiring other ester names or balanced equations.
Explain weak carboxylic acids using partial ionisation in aqueous solution.
Compare pH and hydrogen-ion concentration at the same stated acid concentration.
Recognise alkene monomers and addition polymers from their structures.
Draw a repeating unit with continuation bonds and explain conservation of the monomer atoms.
Relate a supplied alkene monomer to its addition-polymer repeating unit.
Retain side groups and carbon valencies, using propene and a supplied substituted alkene.
Explain condensation polymerisation using monomers with two functional groups.
Identify the ester linkage and relate a polyester repeating unit to diol and dicarboxylic-acid monomers.
Identify amino and carboxylic-acid groups in glycine and explain condensation to polypeptides.
Explain that different amino acids can combine in the same chain to produce proteins.
Describe most DNA molecules as two nucleotide-polymer chains forming a double helix.
Identify four nucleotide types and DNA’s role in encoding genetic instructions.
Name amino acids as protein monomers and glucose as the monomer type for starch and cellulose.
Distinguish common monomer identity from polymer arrangement and biological role.

functional group: An atom group determining characteristic reactions
isomer: A compound sharing a formula but differing in structure
crude oil: A finite mixture of many compounds, mainly hydrocarbons, formed from ancient biomass
hydrocarbon: A compound containing carbon and hydrogen only
alkane: A saturated hydrocarbon in the homologous series with general formula CnH2n+2
displayed formula: A representation showing every atom and each bond in a molecule
fraction: A separated mixture of hydrocarbons with similar carbon numbers and boiling ranges
feedstock: A starting material used to manufacture other chemicals or materials
viscosity: Resistance of a fluid to flow
flammability: How readily a substance ignites under the stated conditions
complete combustion: Burning a hydrocarbon in sufficient oxygen to form carbon dioxide and water
oxygen supply: The oxygen available for a combustion reaction
cracking: Breaking larger hydrocarbon molecules into smaller, more useful molecules
alkene: An unsaturated hydrocarbon containing a carbon–carbon double bond
unsaturated: Containing a carbon–carbon multiple bond that can undergo addition in the stated hydrocarbon model
hydrogenation: Addition of hydrogen across a carbon–carbon multiple bond to form a more saturated product
addition reaction: A reaction adding atoms across a multiple bond without removing the existing carbon skeleton
halogen addition: Addition of one halogen atom to each carbon of a carbon–carbon double bond
decolourisation: Loss of a reagent’s visible colour during the stated reaction
hydration: Addition of water across a multiple bond in the stated alkene reaction
steam: Water in the gaseous state
alcohol · ⁨醇⁩: An organic compound containing an –OH functional group attached to its carbon framework
alcohol oxidation: Reaction of the stated alcohol with an oxidising agent under the given conditions
solubility: How much solute can dissolve in a stated amount of solvent under specified conditions
fermentation: Anaerobic conversion of sugars by yeast to an aqueous ethanol solution and carbon dioxide
anaerobic conditions: Conditions excluding oxygen from the stated biological process
carboxylic acid · ⁨羧酸⁩: An organic acid containing the –COOH functional group
carboxyl group: The –COOH group containing C=O and O–H attached to the same carbon
ester · ⁨酯⁩: An organic compound containing the C(=O)–O–C connection formed here from an acid and alcohol
esterification · ⁨酯化⁩: Reaction of a carboxylic acid with an alcohol to form an ester and water
acid strength: The extent of ionisation of an acid in aqueous solution
partial ionisation: Ionisation of only some acid particles in the stated aqueous equilibrium
monomer: A small molecule that can join with others to form a polymer
addition polymerisation: Joining many unsaturated monomers without forming another small molecule in the stated reaction
repeating unit: The smallest stated structural pattern repeated along the polymer chain
side group: An atom or group attached to the stated polymer backbone rather than continuing that backbone
condensation polymerisation: Joining multifunctional monomers while forming small-molecule by-products in the stated reaction
polyester: A polymer containing repeated ester links in its backbone
amino group: The nitrogen-containing –NH₂ functional group in the stated amino-acid structure
peptide link: The C(=O)–NH link formed between the stated amino-acid residues by condensation
nucleotide: A monomer building block of a nucleic-acid chain, containing a sugar, phosphate and a base
double helix: The two-strand helical arrangement of the stated DNA model
natural polymer: A polymer produced in a biological system
glucose: The simple sugar identified as the monomer type for starch and cellulose in this course

4.8 · Chemical analysis

A separation method based on different distributions between stationary and mobile phases.
Rf is distance travelled by a component divided by distance travelled by the solvent front, both measured from the baseline. Compare under the same conditions; an Rf value alone does not establish identity across different solvents.
Use pencil for the baseline, keep spots above solvent level, mark the solvent front promptly, and run known references alongside unknowns. For ion tests, use clean equipment and separate aliquots to avoid carrying reagents into later tests.
The extent to which a sample consists of one element or compound rather than a mixture.
Compare the measured transition with a reference at the same pressure and with calibrated apparatus. A sharp matching melting point supports purity and identity, while a broadened or shifted interval can indicate an impurity. Boiling temperature depends on pressure and the mixture; do not claim every impurity must shift every boiling point in the same direction. A single matching temperature is supporting evidence rather than unique proof of identity, because different substances can have similar values.
Read the temperature at which melting begins and the temperature at which the last solid disappears. Record an interval rather than selecting its midpoint to hide a broad transition. Compare repeat measurements and instrument resolution. Any actual school melting-point or boiling demonstration uses approved samples and controlled heating; never heat an unknown flammable material or close a heated vessel to obtain a more convenient reading.
A mixture designed as a useful product with components in carefully measured quantities.
Use the supplied information to connect component, quantity and intended property. Too much solvent may make a paint spread easily but reduce coating thickness, while the right binder/pigment balance affects coverage and adhesion. These are product-specific trade-offs rather than a universal recipe. A formulation’s measured composition is different from the fixed chemical proportions inside a compound. Proprietary component names are not required knowledge.
Compare fictional product data using the same amount of painted surface or another stated service. Calculate component quantities for a chosen batch, then check that the total agrees with the target mass. Students can analyse labelled prepared mixtures or a teacher-approved low-risk model; the task is not permission to compound medicines, fuels or cleaning chemicals independently. A useful product claim requires measured performance, not just a plausible ingredients list.
The moving phase carrying components through the stated chromatographic system.
Rf is distance moved by the substance divided by distance moved by the solvent. Measure both from the origin line, using the centre of the spot for the substance. The ratio has no unit and ordinarily lies between zero and one in this model. Compare unknown and reference spots using the same paper, solvent and conditions. A different solvent can change the separation and Rf; it is not a universal identifying number for one substance.
Read each lane separately and count distinct spots. Multiple spots show multiple detectable components. A single spot is consistent with a pure compound, but a mixture may have components that overlap or are not detected in that solvent. Use another suitable solvent or independent evidence to strengthen a purity judgement. Plot measured spot centres clearly and state when broad or streaked spots make the centre uncertain; do not choose a convenient edge.
The marked starting line from which chromatographic distances are measured.
Draw a horizontal pencil origin 2 cm above the lower paper edge. Mark positions away from the edges, place small concentrated spots using separate clean tubes and label in pencil. The handbook aims for 2–3 mm spots and no more than 1 cm water depth, keeping the origin above the water while the paper bottom dips into it. Support the paper without touching the beaker walls and avoid moving the vessel during the run.
Stop before the solvent reaches the paper top, remove the paper and mark the solvent front immediately in pencil. Let the paper dry before measuring from the origin to each spot centre and front. Preserve the actual chromatogram, measurements, solvent identity and calculated Rf. Compare U with references in the same run; record an additional spot as unidentified rather than inventing a name. The school method and risk assessment govern capillary handling and any drying equipment.
A wooden splint with a flame used in the stated gas test.
A reaction prediction and a gas test answer different questions. A known reactive metal with dilute acid can produce hydrogen, but bubbles alone could be another gas in a different system. A positive pop supports hydrogen under the approved test conditions. A weak or absent observation can reflect a small or diluted sample as well as absence of hydrogen, so an inconclusive trial should not be rewritten as a definite negative.
An actual school gas test requires teacher-selected small quantities, approved collection, an open testing path and controlled ignition. Keep any bulk gas-generating vessel and stored flammable material away from the test flame. Never tightly seal an active gas reaction or scale up an explosive demonstration. Record the actual splint condition, sampling method and observed result, with a labelled apparatus sketch; written identification is preparation rather than laboratory completion.
A wooden splint with a glowing tip rather than an existing flame.
Oxygen supports combustion but is not itself described as a flammable fuel in this test. Increased burning is evidence about combustion conditions; the school identification relies on the characteristic relighting observation. Colourless gas or a known source reaction alone is weaker evidence than a correctly performed test. A sample mixed with air or contaminated during collection may give less clear observations, so retain uncertain results rather than inventing an identification.
For actual supervised testing, the teacher chooses the source, collection method and small sample quantity. Prepare the splint with a clearly glowing tip under the approved method, then place it into the collected gas as instructed. Use eye protection and keep ignition arrangements separate from incompatible flammable substances. Record the starting splint state and whether it relights; do not describe the flame’s colour as the required identifying observation.
An aqueous calcium hydroxide solution used in the carbon-dioxide test.
State the reagent and the positive result together. A flame going out is not a unique carbon-dioxide test because several gases fail to support combustion. Under continued excess carbon dioxide the cloudiness can eventually disappear by further reaction; that optional observation does not invalidate the initial specified milky result. Do not substitute that extension for the basic test or make students memorise an unrequired second equation.
Use teacher-approved dilute reagents and an open gas-delivery path that prevents pressure buildup. For a carbonate demonstration, add the selected dilute acid only under the school method, then direct a small gas sample into fresh limewater. Avoid suck-back and unapproved gas generation. Record the initial clarity, reagent identity and actual change, comparing a suitable blank when provided. Distinguish the measured observation from an inferred gas identity.
Removal of the indicator colour so damp litmus becomes white in the stated chlorine test.
Chlorine is Cl₂, a molecular element, while chloride is Cl⁻, a negative ion in compounds and solutions. Their tests are different: gas bleaching identifies chlorine under suitable conditions, whereas acidified silver nitrate precipitates chloride ions. Pale green gas appearance can support the description but does not justify smelling a sample. Not every substance containing chlorine atoms releases chlorine gas or bleaches paper in the same way.
Chlorine is harmful to breathe. Actual testing uses only teacher-approved small quantities, suitable ventilation/containment and controlled handling under the current school risk assessment. Never sniff the gas, generate it from improvised cleaning products or independently repeat an electrolysis setup to obtain it. Students can interpret source observations if direct exposure is unsuitable; the teacher records what actual practical experience and demonstration evidence were provided.
A qualitative test using the characteristic flame colour produced by a stated metal-ion sample.
A sample is introduced into a suitable flame using the school-approved wire or soaked-splint method. Compare its emission with known references under the same conditions. A mixture may have masked colours; intense sodium emission can obscure another response. Yellow therefore supports sodium in the tested conditions but does not prove no other ions are present. Contamination of the wire, sample or apparatus can produce a misleading result.
For actual RP7 school work, use labelled known compounds, approved small amounts and a blue Bunsen flame under supervision. Clean the nichrome wire by the approved method between samples; the acquired handbook suggests fine emery paper and explicitly excludes concentrated hydrochloric-acid watch glasses for GCSE students. Record actual observed colour before naming the cation. Prepared reference photographs or a teacher demonstration can support interpretation but do not silently certify individual practical performance.
The school-selected single ionic compound whose cation and anion are to be identified.
Establish reference results with labelled known cation compounds and carbonate/sulfate/chloride/bromide/iodide salts. Use fresh unknown portions for different tests. Keep flame-wire cleaning, dilute-acid identities and precipitating reagents exact. Record carbonate gas and limewater, sulfate with dilute HCl/barium chloride, and halides with dilute nitric acid/silver nitrate. Compare subtle silver-halide colours side by side. The handbook suggests potassium sulfate as a convenient unknown; the school can choose another suitable single compound.
Use the approved station arrangement, separate pipettes, current school risk assessment and supervised small quantities. The teacher demonstrates pipette gas transfer to limewater and controls Bunsen use, caustic reagents, soluble barium salts and silver waste. Keep raw observations, test conditions, any no-change/inconclusive result and justified inference. Do not erase contradictory results; check contamination, repeat under approved conditions and retain the explanation. No proprietary reagent-making or concentrated-acid cleaning is assigned to students.
An insoluble solid formed in the stated reaction from dissolved substances.
Record initial precipitation separately from what happens on adding excess reagent. A white precipitate that dissolves in excess supports aluminium; one remaining supports calcium or magnesium within the supplied set. Use additional suitable evidence, such as calcium’s orange-red flame, when a single observation leaves alternatives. The Roman numerals in iron(II) and iron(III) distinguish ion charges and their different hydroxides. Colour should be compared with known results in similar conditions.
Carry out only the approved supervised small-scale comparison using labelled solutions, separate clean portions and selected dilute sodium hydroxide. Add the reagent as instructed, observe the precipitate, then add excess and record whether it dissolves. Sodium hydroxide is corrosive at sufficient concentration; school controls govern eye protection, handling and waste. Students need not write an equation for producing sodium aluminate: the specification requires the observation but excludes that equation.
The stated metal-to-hydroxide-group composition of a metal hydroxide.
The divalent examples have form MCl₂(aq) + 2NaOH(aq) → M(OH)₂(s) + 2NaCl(aq), with M the stated Mg, Ca, Cu or Fe(II). For Al and Fe(III), MCl₃(aq) + 3NaOH(aq) → M(OH)₃(s) + 3NaCl(aq). These are full symbol equations, suitable for the both-tier requirement. Separate HT ionic-equation work may simplify them, but it is not needed to recognise the common-tier formation equations here.
Write actual names and formulae first, then count metal, chlorine, sodium, oxygen and hydrogen on both sides. Keep the precipitate solid and the stated starting solutions aqueous. Use the visible precipitate observation to justify (s), rather than assuming every product is solid. Compare the six formation equations with the identification table. The aluminium excess-alkali dissolution observation is separate; the specification does not require a sodium-aluminate production equation.
Addition of dilute acid followed by identifying the evolved carbon dioxide with limewater.
For sodium carbonate and hydrochloric acid, Na₂CO₃ + 2HCl → 2NaCl + CO₂ + H₂O illustrates the source of the gas. The CO₂ test uses aqueous calcium hydroxide and forms insoluble calcium carbonate. Effervescence means gas bubbles, not a unique gas identity. A flame being extinguished is also non-specific: several gases do not support combustion, so it is weaker than the specified limewater result.
RP7’s acquired student method uses a small gas transfer with a teat pipette into limewater. The teacher demonstrates this and several transfers may be needed before cloudiness is clear. Use the school-selected dilute acid and small quantities with eye protection, no mouth pipetting and no sealed gas-producing vessel. Record actual effervescence, transfer conditions and limewater change; a failed transfer should be documented as an observation limit, not renamed as another anion.
The stated dilute-nitric-acid and silver-nitrate precipitation test for halide ions.
Nitric acid does not introduce a halide ion that would contaminate this test. Using hydrochloric acid instead introduces chloride and can produce a misleading white precipitate with silver nitrate. Acidification also removes carbonate interference in the standard comparison. The school method governs quantities and handling. Silver-halide colours are subtle, so compare known results side by side under the same light rather than relying on a tiny colour difference from memory.
Use a fresh clean portion of each solution, selected dilute nitric acid and approved silver nitrate under school supervision. Keep pipettes separate and record observations before assigning identities. Silver salts and acids require the current risk assessment, eye protection and controlled waste collection. Do not mix every reagent into the same unknown tube; previous tests may introduce ions that affect the next test. Additional ammonia-solubility chemistry is not required by this section.
The stated dilute-hydrochloric-acid and barium-chloride test producing a white sulfate precipitate.
Dilute hydrochloric acid helps remove carbonate interference before barium chloride is added. Barium salts can otherwise precipitate with carbonate as well. Do not use sulfuric acid to acidify the sample: it introduces sulfate and could produce a false positive. This requirement differs from the halide test, where nitric acid avoids introducing chloride. The acid choice is therefore part of the chemical evidence, not a interchangeable label.
For actual RP7, take a fresh portion of the solution, add the teacher-selected dilute hydrochloric acid, then the approved barium chloride solution. Use clean pipettes and compare with known reference results. Soluble barium salts require careful handling and controlled waste under the school risk assessment. Record the observation before naming sulfate; avoid contaminating the stock or treating the white material as safe merely because insoluble barium sulfate is the intended product.
The ability to detect small amounts or concentrations in the stated analytical comparison.
Compare an instrument and a chemical test against the same task. Detecting a small metal-ion concentration is different from identifying a high-concentration salt by flame colour. Instrumental outputs can be numerical and easier to compare objectively than a subtle cream/white judgement. Instruments still need appropriate references, calibration, clean samples and correct operation. A fast result can be wrong if contamination or an unsuitable reference invalidates the method.
Use a supplied table of reference concentrations, measured values, minimum detectable concentrations and run times. Link each advantage to the relevant column rather than repeating three adjectives with no evidence. Consider school access, preparation and cost as practical limitations when information is provided. This section asks for advantages over the stated chemical tests; it does not require operating an unacquired advanced instrument or memorising unrelated spectroscopic mechanisms.
An emission output with distinct line positions compared with supplied references.
Separate line position from signal strength. Position supports identity; intensity at a suitable line, compared with known concentrations, supports quantity. Use the supplied wavelength scale or reference table, not a memorised list of invented atomic wavelengths. The original figure’s reference X/Y positions are explicitly fictional teaching data. Check several lines and overlapping signals; one shared line alone can leave alternatives. No electron-transition or advanced instrument mechanism is required here.
For a supplied calibration, plot the measured reference signal against known concentration and read an unknown within the calibrated range. Use the same sample preparation, instrument conditions and background correction described by the data. Avoid extrapolation beyond the references without evidence. A teacher-approved handheld spectroscope observation can show line structure; it does not by itself supply quantitative concentration calibration or replace the RP7 chemical-test experience.
Distinguish a single element or compound from a mixture and from everyday uses of pure.
Interpret melting/boiling data under comparable conditions as evidence of purity.
Identify a formulation as a mixture designed for a useful product.
Explain component functions and measured proportions using supplied product information.
Explain separation through different distributions between stationary and mobile phases.
Interpret chromatograms, calculate Rf and compare references under the same conditions.
Perform school-supervised paper chromatography of known and unknown coloured substances.
Record spot/front distances, calculate Rf and evaluate unknown components against references.
Describe the burning-splint hydrogen test at the open end of a gas test tube.
Distinguish a positive gas observation from an untested source-reaction prediction.
Describe insertion of a glowing splint and its relighting in oxygen.
Use the exact splint condition and observation to interpret supplied gas evidence.
Describe the carbon-dioxide test with aqueous calcium hydroxide.
Interpret milky/cloudy limewater as the specified positive observation.
Describe damp litmus bleaching white in chlorine gas.
Distinguish chlorine gas from chloride ions and preserve approved containment.
Identify lithium, sodium, potassium, calcium and copper ions from the specified flame colours.
Explain masking/contamination limits and compare unknowns with clean references.
Perform supervised reference/unknown ion tests using fresh portions and recorded observations.
Identify cation and anion from the appropriate test matrix while documenting uncertainty.
Identify the six specified cations from sodium-hydroxide precipitates and excess-alkali behaviour.
Use further evidence when white precipitates do not uniquely distinguish calcium and magnesium.
Write balanced symbol equations for the specified insoluble metal hydroxides.
Retain formulae, brackets and state symbols while balancing every element.
Describe dilute-acid effervescence followed by the limewater test for carbon dioxide.
Distinguish the carbonate-ion inference from a non-specific gas observation.
Identify chloride, bromide and iodide from white, cream and yellow silver-halide precipitates.
Specify dilute nitric acid and silver nitrate, with clean fresh sample portions.
Describe dilute hydrochloric acid and barium chloride producing white barium sulfate.
Distinguish sulfate evidence from other tests that also give white precipitates.
State advantages of instrumental detection and identification over the specified qualitative tests.
Distinguish accurate results, detecting small amounts and rapid measurement.
Describe flame emission through a spectroscope to a line-spectrum output.
Interpret supplied reference spectra and concentration data without unrequired atomic mechanisms.

chromatography: A separation method based on different distributions between stationary and mobile phases
Rf: Spot distance divided by solvent-front distance
chemical purity: The extent to which a sample consists of one element or compound rather than a mixture
melting interval: The measured temperature range from the start to the completion of melting in the stated test
formulation: A mixture designed as a useful product with components in carefully measured quantities
component proportion: The stated fraction or percentage of a component in the whole mixture
mobile phase: The moving phase carrying components through the stated chromatographic system
stationary phase · ⁨固定相⁩: The phase that remains in place while components distribute between it and the mobile phase
origin line: The marked starting line from which chromatographic distances are measured
solvent front: The farthest position reached by the moving solvent in the stated chromatogram
burning splint: A wooden splint with a flame used in the stated gas test
positive test: An observed result matching the stated identifying response under suitable test conditions
glowing splint: A wooden splint with a glowing tip rather than an existing flame
relighting: Resumption of a flame from the glowing splint in the stated oxygen test
limewater: An aqueous calcium hydroxide solution used in the carbon-dioxide test
cloudy result: Loss of solution clarity through the stated suspended solid in a positive limewater test
bleaching: Removal of the indicator colour so damp litmus becomes white in the stated chlorine test
damp litmus: Litmus paper moistened for the specified chlorine gas test
flame test: A qualitative test using the characteristic flame colour produced by a stated metal-ion sample
cation: An ion with a positive electrical charge
unknown salt: The school-selected single ionic compound whose cation and anion are to be identified
reference matrix: A recorded set of known reagent-and-observation results used to interpret the unknown
precipitate: An insoluble solid formed in the stated reaction from dissolved substances
excess reagent: More reagent than needed for the stated initial reaction, allowing subsequent behaviour to be observed
hydroxide formula: The stated metal-to-hydroxide-group composition of a metal hydroxide
formation equation: A balanced equation describing production of the stated hydroxide precipitate
carbonate test: Addition of dilute acid followed by identifying the evolved carbon dioxide with limewater
effervescence: Formation of gas bubbles observed in the stated reaction
halide test: The stated dilute-nitric-acid and silver-nitrate precipitation test for halide ions
silver halide: The insoluble silver compound formed with the stated chloride, bromide or iodide ion
sulfate test · ⁨硫酸盐测试⁩: The stated dilute-hydrochloric-acid and barium-chloride test producing a white sulfate precipitate
test interference: Another species or contamination producing a misleading response in the stated test
analytical sensitivity: The ability to detect small amounts or concentrations in the stated analytical comparison
detection threshold: The lowest stated amount or concentration giving a detectable response under the supplied conditions
line spectrum: An emission output with distinct line positions compared with supplied references
calibration data: Measurements from known concentrations used to relate the stated signal to an unknown concentration

4.9 · Chemistry of the atmosphere

Assessment across production, use and disposal.
Define the functional unit before comparing products. The same delivered service, such as carrying one litre of water a hundred times, is fairer than comparing one object with another regardless of lifetime.
List system boundaries, energy sources and assumptions. Compare water demand, emissions and waste separately before making a judgement. Explain whose priorities affect the decision and where the data are uncertain.

life-cycle assessment: Assessment across production, use and disposal
functional unit: The common service used for a fair comparison

4.10 · Using resources

Assessment across production, use and disposal.
Define the functional unit before comparing products. The same delivered service, such as carrying one litre of water a hundred times, is fairer than comparing one object with another regardless of lifetime.
List system boundaries, energy sources and assumptions. Compare water demand, emissions and waste separately before making a judgement. Explain whose priorities affect the decision and where the data are uncertain.

life-cycle assessment: Assessment across production, use and disposal
functional unit: The common service used for a fair comparison

Preparing for this qualification · ⁨备考指南⁩

Foundation and Higher routes must retain Chemistry-only and HT-only statements.
Paper 1: 4.1–4.5; Paper 2: 4.6–4.10. Each paper is 100 marks, 1 h 45 min, 50%.
Moles, quantitative titration and equilibrium depth are gated by tier; practical experience is not a substitute written assessment.

Teaching coverage still needed · ⁨仍需教学覆盖内容⁩

Teaching reviewed for all 15 substantive sections 4.1.1–4.1.3, with separate HT half/ionic equations and both-tier Chemistry-only transition metals. Whole-unit assessments, complete bank review and qualification mocks remain unfinished.
Exact teaching reviewed for all 18 substantive sections 4.2.1–4.2.4, including separate HT particle-model evaluation and both-tier Chemistry-only nanoparticles. Whole-unit tests, full bank review and qualification mocks remain unfinished.
Exact 4.3.1.1–4.3.5 teaching reviewed: shared conservation/formula-mass/uncertainty/mass-concentration and Chemistry-only yield/atom economy stay distinct from HT mole, embedded explanation, theoretical-yield, pathway, molar-concentration, titration and gas-volume work. Whole-unit assessments, qualification mocks and full-paper review remain unfinished.
Exact 4.4.1.1–4.4.3.5 teaching reviewed: room-temperature reactivity, extraction, acid/base/salt chemistry and actual RP1–3. Electron-redox, RP2 concentration, strong/weak acid and electrode half equations have explicit HT cases. Whole-unit tests, complete bank review and mocks remain unfinished.
Exact 4.5.1.1–4.5.2.2 teaching reviewed: temperature-change evidence and actual RP4, curved profiles and activation barriers, separate Chemistry cells/fuel-cell evaluation, and distinct HT bond-energy/half-equation cases. Solution energy/ΔH calculations are not required. Whole-unit tests, full bank review and mocks remain unfinished.
Exact 4.6.1.1–4.6.2.7 teaching reviewed: common-tier mass/volume rates and tangent interpretation, both actual RP5 methods, collision factors, catalysts, reversible reactions and dynamic equilibrium. Numerical tangent gradients/mol/s and the four equilibrium condition sections remain explicitly Higher-only. Whole-unit tests, full-paper review and mocks remain unfinished.
Exact 4.7.1.1–4.7.3.4 teaching reviewed: crude oil/fractions/trends/combustion/cracking, first-four displayed alkenes and all specified addition products/conditions, alcohol structures/uses/reactions/fermentation, acid reactions and ethyl ethanoate, with separate HT weak-acid explanation. Addition-polymer monomer/repeat drawing, HT condensation/polyester and amino-acid chemistry, and both-tier DNA/protein/starch/cellulose monomers reviewed. Whole-unit assessment and bank parity remain unfinished.
Exact 4.8.1.1–4.8.2.4 purity, formulations, phase distribution/Rf and all four specified gas tests reviewed, with actual RP6 chromatography evidence. All 4.8.3 ion tests, balanced hydroxide formation, instrumental advantages, supplied flame spectra/calibration and actual RP7 evidence reviewed. Whole-unit assessments and full bank review remain unfinished.
Atmospheric evolution, carbon dioxide/methane climate evidence and pollutants need full objective coverage.
Potable water, wastewater, alternative metal extraction, corrosion, alloys, composites and Haber/NPK chemistry need full coverage.

Specifications and sample documents · ⁨课程大纲和样件文件⁩

Course materials · ⁨课程资料⁩

Course preparation · ⁨课程准备⁩

Documents are available. Board-specific notes, assessments and interactive past-paper practice are not yet available for every course. · ⁨文档已提供。并非所有课程都具备考试局特定的注释、测评及交互式历年真题练习。⁩

Lessons · ⁨课程⁩ →