Edexcel GCSE Chemistry Revision

Explore Edexcel GCSE Chemistry revision topics. Check the official specification for complete requirements for your exam year and route.

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Edexcel · GCSE Chemistry1CH0Exam years · 2027

Assessment The assessment includes Paper 1 and Paper 2; see the official specification for the full structure.

Official Edexcel specification (opens in a new tab)

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Open a point to read what to know, how to use it in an assessment, a plain-language explanation and a self-check. Answers stay hidden until you choose to reveal them.

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Paper tier

Foundation tier
Higher tier

§3, pp. 41–42; 1CH0/1F and /2F or 1CH0/1H and /2H; all assessments in the qualification must use one tier

Navigate the two 1CH0 papers and tier routeCore
Scope: Core · §3, pp. 41–43; Paper 1/2 structure, tiers, overlap and Assessment Objectives
Open revision guidance
What to know
Separate Chemistry 1CH0 has two 100-mark, 1 hour 45 minute papers. Paper 1 assesses Topics 1–5; Paper 2 assesses Topic 1 and Topics 6–9, so Topic 1 knowledge can recur on either paper. Both papers count for 50% of the GCSE.
Choose one route: Foundation 1CH0/1F and /2F target grades 1–5; Higher 1CH0/1H and /2H target grades 4–9. A student must take both papers in the same tier. Each paper contains 27 marks of questions shared across tiers. Appendix 8 distinguishes the RQF qualification number 601/8611/2, GCSE subject code 1CH0 and individual paper codes; the two entered paper codes must match the chosen tier.
Across the GCSE AO1 (knowledge/understanding of scientific ideas, techniques and procedures) is 40% ±3 percentage points; AO2 (application of scientific ideas and enquiry/techniques/procedures) is 40% ±3; AO3 (analysis/evaluation, judgement/conclusions and improving procedures) is 20% ±3. Each paper contributes 20% AO1, 20% AO2 and 10% AO3 to the GCSE. Papers include multiple choice, short-answer, calculations and extended open-response; calculators may be used.
The specification describes all assessment as in May/June of the year students sit the qualification; Pearson's final Summer 2027 GCSE timetable schedules both 1CH0 Chemistry papers in each tier. Appendix 9 permits a desk-suitable battery/solar calculator without lids/cases/covers carrying instructions/formulae; it must not translate languages, perform symbolic algebra/differentiation/integration, communicate or access the internet, or store retrievable information (including databanks, dictionaries, formulae or text). Candidates are responsible for power, working condition and clearing stored data; calculators cannot be borrowed from another candidate.
What to do in an assessment
Check the paper code (1CH0/1F or /1H; 1CH0/2F or /2H), selected tier and command word. Both papers in one qualification must use the same tier and are sat in the same exam series/year.
Practise recall, application, calculations, practical-method questions, data interpretation and extended explanations; show equations/substitution/unit conversions for calculations and support evaluations with evidence.
In plain language
You sit two papers at one tier. They cover separate topic groupings but share Topic 1 and both test practical, mathematical and scientific reasoning.
Self-check
Can a student enter Paper 1 Foundation and Paper 2 Higher?
Show the answer
No. Both assessments for 1CH0 must be completed at the same tier.
Use chemical formulae, equations and risk controlsCore
Scope: Core · Topic 0, points 0.1–0.6
Open revision guidance
What to know
Recall common element formulae H₂, N₂, O₂, F₂, Cl₂, Br₂ and I₂ (noble gases are monatomic); common compounds H₂O, CO₂, CO, CH₄, NH₃, HCl, H₂SO₄, HNO₃ and NaOH; common ions H⁺, OH⁻, NH₄⁺, Na⁺, K⁺, Ag⁺, Ca²⁺, Mg²⁺, Zn²⁺, Cu²⁺, Fe²⁺, Fe³⁺, Al³⁺, Cl⁻, Br⁻, I⁻, O²⁻, CO₃²⁻, SO₄²⁻ and NO₃⁻. An element formula describes its particles, a compound formula gives particle composition and an ionic formula balances positive/negative charge; retrieve formula, name and charge from both directions.
Translate a reaction into a word equation and then a balanced symbol equation. Balance by changing coefficients, never subscripts; include (s), (l), (g) or (aq) where known. Example: magnesium ribbon + hydrochloric acid → magnesium chloride + hydrogen; Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g). Count each atom on both sides.
Recognise the common laboratory pictograms: flame (flammable), flame over circle (oxidising), exploding bomb (explosive), gas cylinder (gas under pressure), corrosion (skin/eye burns or metal corrosion), skull (acute toxicity), exclamation mark (irritant/harmful), health-hazard silhouette (serious/chronic health effects), and dead tree/fish (environmental hazard). Pictograms communicate specific danger and guide precautions; risk assessment also considers exposure and procedure. Examples across this course: dilute HCl/H₂SO₄/NaOH or BaCl₂ splashes require goggles and immediate rinsing access; flammable alcohols and paraffin vapour stay away from ignition; avoid inhaling sulfur dioxide from thiosulfate or powdered copper oxide; handle Bunsen-heated glassware/crucibles with tongs and allow cooling; do not touch electrodes until the DC supply is switched off; use ventilation for chlorine/ammonia tests and dispose of metal-ion/precipitation waste as instructed.
What to do in an assessment
Check atom counts on each side of a balanced symbol equation; keep each formula fixed and change coefficients only. Write state symbols when the state is known.
For each practical risk, name the reagent/apparatus hazard, possible exposure, consequence and a proportionate control; check controls against the procedure, not the pictogram alone.
In plain language
Chemical formulae and equations are compact descriptions of particles and reactions; hazard labels help you choose safe procedures.
Self-check
Balance magnesium + hydrochloric acid → magnesium chloride + hydrogen, including states: Mg(s) + HCl(aq) → MgCl₂(aq) + H₂(g).
Show the answer
Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g). There is one Mg, two Cl and two H atoms on each side; formula subscripts are unchanged.
Connect atomic structure, bonding and quantitative chemistryCore
Scope: Core · Topic 1, points 1.1–1.53
Open revision guidance
What to know
Dalton proposed solid, indivisible atoms; discovery of the electron showed atoms contain smaller negative particles, scattering evidence led to a tiny positive nucleus, and later evidence established neutrons and electrons in shells. An atom is mostly empty space; nearly all mass is in the very small nucleus. Proton charge/mass are +1/1, neutron 0/1 and electron −1/negligible. Neutral atoms have equal proton/electron counts. Mass number = protons + neutrons; atomic number = proton count. Isotopes share proton number but differ in neutrons; this makes some relative atomic masses non-integers. Atomic diameter is about 10⁻¹⁰ m (0.1 nm); a small molecule is roughly 10⁻⁹ m. Thus one molecule is around ten atomic diameters across, while a 1 mm grain is about 10⁷ times an atom's diameter. Estimate and compare sizes by powers of ten, not to misleading drawn scale.
Mendeleev arranged known elements using properties of elements and compounds, left gaps and predicted undiscovered elements. The modern table is ordered by atomic number (proton number), not always relative atomic mass; periods are rows and groups are columns with similar properties. First-20 electron configurations 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; K 2,8,8,1; Ca 2,8,8,2. Draw shells with the nucleus labelled; outer-electron number identifies main-group column and occupied-shell number the period. Metals lie mainly left/centre with few outer electrons; non-metals mainly right with nearly full outer shells.
An ion is a charged atom or group. Form cations by electron loss and anions by gain; in simple ions electrons = proton number adjusted for charge. Worked common-tier counts: ²⁴Mg²⁺ has 12 protons, 12 neutrons and 10 electrons (24−12; 12−2); ³⁵Cl⁻ has 17 protons, 18 neutrons and 18 electrons (35−17; 17+1). Dot-and-cross diagrams show outer electrons transferred, not the later lattice. For groups 1, 2, 6 and 7, use charges +1, +2, −2 and −1. Balance charge to deduce common ionic formulae: Na⁺/O²⁻ → Na₂O; Mg²⁺/OH⁻ → Mg(OH)₂; Al³⁺/Cl⁻ → AlCl₃; Ca²⁺/NO₃⁻ → Ca(NO₃)₂; Mg²⁺/CO₃²⁻ → MgCO₃; Al³⁺/SO₄²⁻ → Al₂(SO₄)₃. These demonstrate oxide, hydroxide, halide, nitrate, carbonate and sulfate; use brackets when more than one polyatomic ion is needed. Use -ide for binary compounds and -ate where the ion contains oxygen.
Use the chemistry-diagrams visual to trace sodium's transferred outer electron into chloride and to compare the covalent shared pairs and lone pairs in H₂, HCl, H₂O, CH₄, O₂ and CO₂. Dots/crosses identify which atom contributed each electron; count the outer-shell electrons shown and distinguish ionic transfer from covalent sharing.
Giant ionic solids are regular lattices of alternating charges. Translate a 3-D lattice/ball-and-stick view into a 2-D diagram and back; a flat diagram cannot show full depth or every surrounding ion. Ball-and-stick models distort particle size/spacing.
Classify ionic, simple molecular, giant covalent and metallic structures. Compare melting/boiling point, water solubility and conductivity in solid/molten/solution states: ionic NaCl dissolves and conducts only molten/aqueous; simple molecular sugar dissolves but its solution does not conduct; giant covalent silicon dioxide has high melting point and is insoluble/non-conducting; metals such as copper conduct as solids and are generally insoluble. These are comparisons, not universal solubility rules. Metals are generally shiny, solid (except mercury), high-melting/high-density and good electrical conductors; non-metals are usually low-boiling and poor conductors (with important exceptions such as graphite). Diamond has a 3-D network of four strong covalent bonds per C, so it is very hard and used for cutting tools. Graphite has three bonds per C in layers, delocalised electrons for electrical conduction/electrodes and weak interlayer forces for lubrication. C₆₀ is a hollow fullerene cage of pentagons/hexagons; graphene is one hexagonal carbon layer, strong and conducting. Their structures explain distinct properties/uses. Metals have positive ions in delocalised electrons; layers can slide while attraction remains, explaining malleability and conductivity. Poly(ethene) is a long carbon-chain polymer.
Calculate Mr by summing relative atomic masses. Percentage by mass = element mass contribution in formula ÷ Mr × 100%. Empirical formula: turn masses/percentages into moles, divide by smallest and scale to whole numbers; molecular formula is an integer multiple consistent with relative molecular mass. Worked common-tier conversion: C₄H₈ has C:H atom ratio 4:8; divide both by 4 to obtain empirical formula CH₂. Conversely, if the empirical formula is CH₂ and Mr=56, its multiplier is 56/14=4, so the molecular formula is C₄H₈. Another worked mass-to-moles example: 24 g carbon and 4 g hydrogen give 2 mol C and 4 mol H, ratio CH₂. For MgO empirical-formula practical, weigh clean crucible+lid, add magnesium and reweigh; heat with lid slightly ajar to admit oxygen but reduce loss, cool with lid and reweigh crucible+product. Repeat heat/cool/weigh to constant mass. Subtract crucible mass for product mass, subtract magnesium mass for oxygen mass, divide each element mass by Ar and reduce mole ratio.
Conservation of mass: a closed flask containing a precipitation reaction has equal total mass before and after because no matter enters/leaves. In an open flask the measured mass can change if gas leaves or enters; atoms are still conserved. For 2Mg + O₂ → 2MgO, 48 g Mg has 2 mol, so forms 2 mol MgO = 80 g (using Mg=24, O=16). For concentration in g dm⁻³, c = mass of dissolved solute (g) ÷ solution volume (dm³); 5.0 g in 0.250 dm³ gives 20 g dm⁻³. Convert cm³ to dm³ by dividing by 1000.
What to do in an assessment
For atom/ion questions, use atomic number for protons, mass number minus atomic number for neutrons, then adjust proton number by the signed ionic charge to find electrons (subtract for a positive charge; add for a negative charge). Draw shell diagrams and check total electrons.
For formula-mass or concentration work, write the relationship, substitute values with units, rearrange if needed and report a sensible significant figure. For formulae, balance ionic charge to zero, including polyatomic ions and brackets where required.
Explain a material property as structure → bond/attraction → particle movement or energy needed; compare both named substances using the property asked.
In plain language
Atomic number and electron arrangement explain the periodic table; bonding and structure explain material properties, and arithmetic links formulae to measured quantities.
Self-check
An atom has atomic number 11 and mass number 23. How many protons, neutrons and electrons does a neutral atom have?
Show the answer
11 protons, 23 − 11 = 12 neutrons and 11 electrons.
View diagrams
Sodium chloride: ionic transfer
Before transfer, sodium is 2,8,1 and chlorine is 2,8,7. The sodium outer electron transfers to chlorine. Separate bracketed ions result: Na⁺ has configuration 2,8 and Cl⁻ has 2,8,8. Dots identify sodium-origin electrons; crosses identify chlorine-origin electrons. This is electron transfer, not a shared covalent pair.
Hydrogen · H₂
Two hydrogen atoms contribute one electron each to one shared pair. The dot comes from the left H atom and the cross from the right H atom; each hydrogen now has a full first shell.
Hydrogen chloride · HCl
One H electron and one Cl electron form one shared pair. Chlorine contributes six more electrons as three lone pairs; all seven of its outer-shell electrons are shown.
Water · H₂O
Oxygen shares one electron with each H, making two shared pairs. Its four remaining outer-shell electrons make two lone pairs.
Methane · CH₄
Carbon shares one electron with each of four H atoms, making four shared pairs. All four carbon outer-shell electrons and all four hydrogen electrons are shown; there are no lone pairs.
Oxygen · O₂
The two oxygen atoms share two pairs, forming a double bond. Each oxygen keeps two lone pairs: 12 outer-shell electrons are shown altogether.
Carbon dioxide · CO₂
Carbon forms a double bond to each oxygen: four shared pairs in total. Each oxygen has two lone pairs; carbon has none. All 16 outer-shell electrons are shown.
In dot-and-cross diagrams, a dot is an electron from the atom named in the explanation and a cross is from the other atom. The symbols show electron origin, not a charge. Shared pairs belong between bonded atoms; lone pairs stay on one atom.
The dot-and-cross and ionic shell sketches show outer-shell electrons only; inner-shell electrons are omitted from those sketches. Full sodium and chlorine configurations are printed beside the ions. Displayed structural formulae instead draw every atom, including each hydrogen, and every covalent bond.
Trace the Na-to-Cl electron transfer and distinguish it from the covalent shared pairs in the six molecule panels. Use the dot/cross marks to identify electron origin and count shared and lone pairs; the shell sketches show outer-shell electrons.
Explain states and separate mixturesCore
Scope: Core · Topic 2, points 2.1–2.12; core practical 2.11
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What to know
In a solid, particles vibrate in fixed positions and have the lowest average particle energy; liquid particles remain close but move past one another and have greater average energy; gas particles move rapidly/randomly far apart and have the greatest average energy. State changes alter particle energy and arrangement, not chemical identity. A physical mixture can be separated without making new substances.
Melting is solid→liquid, freezing liquid→solid, boiling/evaporation liquid→gas, condensation gas→liquid, sublimation solid→gas and deposition gas→solid. During heating, energy changes particle kinetic energy between state changes and overcomes attractions during a state change; particles do not change chemical identity. Use supplied melting/boiling-point data to decide state: below melting point solid, between melting and boiling liquid, above boiling point gas.
In chemistry, pure means one substance only; everyday 'pure' may mean uncontaminated or natural. A pure substance has a sharp melting/boiling point, whereas impurities usually broaden/depress the melting range. If a sample melts over a wide, lower-temperature range than a pure reference, infer impurity; use the supplied reference/data and apparatus uncertainty rather than claiming an exact contaminant. Choose separation by the property that differs: filter sand from water; evaporate/crystallise dissolved salt; simple-distil salt water and collect water; fractional-distil two miscible liquids with close boiling points in a temperature-gradient column. For an unfamiliar mixture, list components and differing solubility/particle size/boiling point, then sequence the required methods (e.g. filter sand from salt solution, then evaporate filtrate to recover salt).
Chromatography separates soluble substances because they partition differently between a mobile solvent and stationary paper. Pencil (not ink) baseline remains above solvent. Spot and dry samples, place paper in covered solvent chamber, allow solvent to rise, remove and immediately mark solvent front; dry and measure from baseline. Rf = distance travelled by substance ÷ distance travelled by solvent front. Rf has no units and comparisons require same solvent/paper/conditions. Multiple spots indicate an impure mixture; compare spots and Rf with reference substances.
Core practical 2.11 has two parts: chromatography separates dye components; simple distillation recovers pure water as the distillate, leaving dye in the flask. Set-up diagram in words: Bunsen under clamped round-bottom flask (ink + anti-bumping granules) → still head/thermometer → sloping Liebig condenser → open receiver; cooling water enters the condenser's lower side-arm and exits the upper arm. Heat gently, keep condenser water flowing, never seal the receiver, and do not distil to dryness. For chromatography draw pencil baseline above solvent, spot/dry ink, stand paper with solvent below spots in a covered beaker, develop, remove and immediately mark solvent front. Keep solvent away from flame/skin, avoid vapour and do not let spot sit below solvent. Record labelled apparatus, observations and evidence.
Potable water is safe to drink, not chemically pure. Waste/ground water treatment uses sedimentation to settle solids, filtration to remove suspended particles and chlorination to kill microorganisms. Distillation can make seawater potable by evaporating and condensing water, leaving salts behind; water used in analysis must be distilled/deionised so dissolved salts do not interfere.
What to do in an assessment
Choose a method by the physical property that differs; state apparatus, what is collected and why.
Measure chromatography distances from the baseline, use the direction of solvent travel and calculate a unitless Rf.
In plain language
Mixtures can be separated because their components differ in solubility, particle size, boiling point or attraction to a surface.
Self-check
A dye spot moves 2.4 cm while the solvent front moves 6.0 cm. Find Rf.
Show the answer
Rf = 2.4/6.0 = 0.40, with no unit.
Use ions, reactions and electrolysis to explain chemical changeCore
Scope: Core · Topic 3, points 3.1–3.31; core practicals 3.6, 3.17, 3.31
Open revision guidance
What to know
Acids in water supply H⁺ and alkalis supply OH⁻; pH 7 is neutral, below 7 acidic and above 7 alkaline. Litmus turns red in acid and blue in alkali; methyl orange is red in acid/yellow in alkali; phenolphthalein is colourless in acid and pink in alkali.
A base reacts with acid to make salt and water only; an alkali is a soluble base. Metal + acid → salt + hydrogen; oxide/hydroxide + acid → salt + water; carbonate + acid → salt + water + carbon dioxide. Hydrogen gives a squeaky pop with a lit splint; carbon dioxide turns limewater milky.
Neutralisation is H⁺ + OH⁻ → H₂O. For soluble acid and alkali, titrate to the indicator end-point and record exact volumes; repeat with those volumes but no indicator to obtain uncontaminated salt solution, then evaporate/concentrate, crystallise, filter and dry salt. For insoluble reactant, warm acid, add excess solid to consume all acid, filter excess, evaporate some water, cool to crystallise, filter and dry crystals.
Solubility rules: common sodium, potassium and ammonium salts and all nitrates are soluble; common chlorides are soluble except silver/lead chlorides; common sulfates are soluble except lead, barium and calcium sulfates; common carbonates and hydroxides are insoluble except sodium, potassium and ammonium salts. Example: BaCl₂(aq)+Na₂SO₄(aq) forms insoluble BaSO₄(s) and NaCl(aq); AgNO₃(aq)+NaCl(aq) forms AgCl(s) and NaNO₃(aq). Swap ions, apply rules, write formulae and balanced equation, and name the precipitate. Filter through folded paper in a funnel, rinse residue with distilled water to remove soluble ions, then dry on paper/watch glass or in a warm place; washing removes dissolved contaminants and drying gives solid product.
Core practical 3.6: add weighed portions of powdered calcium oxide or hydroxide to one fixed volume of dilute hydrochloric acid, stir and measure pH after every addition with a pH meter/probe or indicator paper on a glass rod. Tabulate cumulative mass added and pH, plot pH (vertical) against mass added (horizontal); keep acid volume/concentration, temperature and stirring consistent. Wear goggles. Core practical 3.17: warm dilute sulfuric acid in a water bath, add copper oxide until excess, filter hot mixture, evaporate blue filtrate in an evaporating basin over Bunsen to near saturation, cool undisturbed to crystallise, filter crystals and dry between filter papers/on a watch glass. Do not heat to dryness; use tongs and goggles.
An electrolyte is an ionic compound molten or dissolved in water; direct current decomposes it. Positive ions move to the negative cathode and negative ions to the positive anode. At a cathode in aqueous solution, H⁺/hydrogen from water competes with metal ions: copper (less reactive than hydrogen) is deposited from CuCl₂ or CuSO₄, whereas sodium does not plate out from NaCl and hydrogen forms. At anode, concentrated halide ions form halogen; in Na₂SO₄/acidified water hydroxide/water gives oxygen. Thus inert-electrode CuCl₂ gives copper and chlorine; concentrated NaCl gives hydrogen and chlorine with NaOH remaining; Na₂SO₄ and acidified water give hydrogen and oxygen; molten PbBr₂ gives lead and bromine. Hydrogen gives a squeaky pop, oxygen relights a glowing splint, chlorine bleaches damp litmus. Predict using competing ions, concentration and electrode material.
In copper sulfate electrolysis, copper appears at the cathode. With copper electrodes, copper enters solution from one electrode and deposits on the other, allowing copper purification; with inert electrodes oxygen is produced at the anode. The copper-sulfate practical compares inert and copper electrodes and records visible products/mass changes.
Core practical 3.31: clean, dry and weigh each copper electrode; set up CuSO₄, set current with DC supply/ammeter and measure electrode masses after fixed electrolysis time. Repeat at different currents, tabulate signed mass changes and plot mass change against current. In a separate run use graphite electrodes and record products. Keep solution concentration/volume, electrode area, spacing and time controlled when varying current; switch off before handling and wear goggles. Core practical 3.6 graphs pH against cumulative mass; practical 3.17 uses water-bath warming, filtration and final crystal drying.
What to do in an assessment
Predict electrode products by listing ions present, distinguishing molten compounds from aqueous solutions and identifying the electrode material.
For a precipitation reaction, apply the solubility rules to the named solutions, name the precipitate, then filter, wash with distilled water and dry it.
In practicals, record readings/observations with units, plot the specified variables and explain controls and safe handling.
In plain language
Acids and bases make salts through neutralisation; ion movement and electrode choice determine products during electrolysis.
Self-check
What is the gas test for carbon dioxide?
Show the answer
Bubble the gas through limewater; it turns milky/cloudy.
Explain metal extraction and reversible equilibriaCore
Scope: Core · Topic 4, points 4.1–4.17
Open revision guidance
What to know
Compare a metal's reaction with water, dilute acid or a salt solution to infer relative reactivity. A more reactive metal displaces a less reactive metal from its salt: magnesium + copper sulfate → magnesium sulfate + copper. Oxygen gain is oxidation and oxygen loss is reduction.
Reactivity series: potassium, sodium, calcium, magnesium, aluminium, carbon, zinc, iron, hydrogen, copper, silver, gold. Metals above hydrogen react with dilute acids to release hydrogen; reaction vigor provides comparisons. Most metals occur in ores in Earth's crust; unreactive metals may be found uncombined. Ores are reduced to extract metal. Metals below carbon can often be obtained by heating their oxides with carbon (e.g. iron oxide); aluminium is too reactive for carbon reduction and requires electrolysis. Compare energy, cost and environmental impacts; blast-furnace detail is not required.
A metal's resistance to oxidation generally increases as reactivity decreases; corrosion is oxidation of a metal. Recycling reduces ore extraction and can conserve energy/material supply, but collection, sorting and processing have costs and impacts.
A life-cycle assessment evaluates environmental effects at raw-material extraction, manufacture, use and disposal. Compare like-for-like data across every stage; identify assumptions, boundaries and trade-offs, and support a decision with evidence rather than a single impact figure.
Worked LCA method: when comparing two materials for the same product, check the functional unit and boundary (e.g. one component over its service lifetime), convert per-kg data to the mass used, total extraction/manufacture/use/disposal impacts, and include replacement lifetime or recycling assumptions. If one material has lower manufacture impact but needs more replacements, compare totals over equal service; state the largest data-supported trade-off and uncertainty before judging. This also applies when comparing virgin metal with recycled metal: include collection/sorting/reprocessing, not only avoided ore extraction.
A reversible reaction is shown with ⇌; changing conditions can alter the direction favoured. Dynamic equilibrium requires a closed system and equal forward/reverse rates; both reactions continue while concentrations remain constant. The Haber reaction is N₂ + 3H₂ ⇌ 2NH₃, using nitrogen from air and hydrogen obtained from natural gas. Conditions are 450 °C, 200 atmospheres and an iron catalyst.
What to do in an assessment
Use observed reactions to place metals in a relative reactivity order; use carbon's position to select a plausible extraction method.
For a life-cycle comparison, state functional unit and boundary, total comparable stages, consider service life and assumptions, then give a supported judgement from the supplied data rather than assuming recycling is always best.
Describe a dynamic equilibrium by closed-system conditions, equal rates and stable concentrations.
In plain language
A metal's reactivity affects how it is extracted. Reversible reactions continue in both directions at equilibrium.
Self-check
Why is aluminium extracted by electrolysis rather than reduction with carbon?
Show the answer
Aluminium is more reactive than carbon, so carbon cannot reduce aluminium oxide; electrolysis is required.
Apply separate-chemistry calculations and industrial ideasCore
Scope: Core · Topic 5, points 5.1C–5.27C
Open revision guidance
What to know
Iron is a transition metal example: it has high melting point (about 1538°C) and density (about 7.9 g cm⁻³), forms coloured compounds (e.g. green Fe²⁺ and brown Fe³⁺ precipitates) and iron is used as a catalyst in Haber ammonia manufacture. Rusting is oxidation and requires both oxygen and water. Prevent oxygen contact with paint/plastic/oil; prevent water contact with grease or dry storage; sacrificial protection attaches a more reactive metal such as zinc, which corrodes in preference to iron.
Electroplating setup: connect object to negative terminal as cathode; place coating metal (e.g. copper) on positive terminal as anode; immerse both without touching in an electrolyte containing coating-metal ions (e.g. copper sulfate for copper plating) and pass direct current. The coating improves appearance/corrosion resistance. Alloy atoms of different sizes distort regular metal layers and impede sliding. Carbon makes iron harder/stronger in steel; chromium/nickel in stainless steel improve corrosion resistance. Aluminium has low density and a protective oxide (aircraft frames); copper conducts (wires); gold is unreactive/malleable (contacts/jewellery); magnalium is light/strong (light components); brass is harder than copper (fittings). Choose the alloy/material for the demanded properties, not its name.
Core practical 5.9C is required at both tiers: rinse pipette with the unknown-concentration sodium hydroxide sample; transfer a fixed accurate aliquot to conical flask, add suitable indicator and place on white tile. Rinse/fill burette with standard HCl, record initial meniscus; add acid while swirling, then dropwise near endpoint. Record final reading and titre; repeat rough then accurate runs to concordant titres and average only concordant results. Read at eye level. Foundation learners must know the safe accurate technique, endpoint, readings, titre and concordant results. The titration mole-ratio arithmetic to calculate an unknown concentration or required volume is the separate Higher-only 5.10C supplement, not a Foundation requirement. Wear goggles; manage corrosive splashes.
Percentage yield = actual yield ÷ theoretical yield × 100%. Example: 18 g actual from 24 g theoretical gives 75%. Actual yield is often lower due to incomplete reaction, product lost during separation/transfer or competing side reactions. Atom economy = Mr desired product ÷ total Mr of all products in balanced equation × 100%. Worked example: CaCO₃→CaO+CO₂; if CaO is desired, Mr(CaO)=56 and total product Mr=56+44=100, so atom economy=56%. A higher value puts a larger fraction of reactant atoms in useful product.
Fertilisers may supply nitrogen, phosphorus and potassium compounds. Ammonia + nitric acid → ammonium nitrate. Laboratory ammonium sulfate: titrate ammonia solution against dilute sulfuric acid, then evaporate/concentrate and crystallise the neutral salt. Industrial production uses multiple stages to make ammonia and sulfuric acid from raw materials and produces much larger quantities; compare the small-scale titration/crystallisation with industrial feedstocks, stages, throughput and control (details of sulfuric-acid manufacture are not required). A chemical cell gives voltage until one reactant is used up. A hydrogen–oxygen fuel cell uses hydrogen and oxygen, producing water as the only product; for a bus operating in a city, compare quiet/zero exhaust at use with hydrogen production, storage, infrastructure, cost and safety.
What to do in an assessment
For the both-tier 5.9C practical, use and read pipette/burette accurately, reach the indicator endpoint, record titre values and identify concordant repeats. Foundation teaching stops at method, endpoint, readings and results; titration mole-ratio calculations for an unknown concentration or required volume belong only to the selected Higher 5.10C supplement.
For yield, show theoretical-yield and actual-yield quantities. For atom economy, name the desired product and use all products in the balanced equation denominator.
Evaluate a material or fuel-cell use by linking each advantage/limitation to its context and supporting a decision with stated evidence.
In plain language
Separate Chemistry adds quantitative tools and applications: titrate accurately, predict product amounts and evaluate how chemical processes make useful substances.
Self-check
A reaction gives 18 g product when the theoretical yield is 24 g. Calculate the percentage yield.
Show the answer
Percentage yield = (18 ÷ 24) × 100 = 75%.
Predict trends in Groups 1, 7 and 0Core
Scope: Core · Topic 6, points 6.1–6.16
Open revision guidance
What to know
Group 1 alkali metals occupy Group 1 and have one outer electron; Group 7 halogens have seven; Group 0 noble gases have full outer shells. Lithium, sodium and potassium are soft with relatively low melting points; each floats/fizzes with water, producing hydrogen and a soluble alkaline hydroxide, more vigorously Li < Na < K. Equations: 2Li+2H₂O→2LiOH+H₂; 2Na+2H₂O→2NaOH+H₂; 2K+2H₂O→2KOH+H₂. Use observations and equations to predict heavier alkali-metal reactions; down the group the outer electron is farther from the nucleus and more shielded, so it is lost more readily.
At room temperature chlorine is a pale-green gas, bromine a red-brown liquid and iodine a grey solid. Down Group 7, colour darkens and melting/boiling points rise. Approximate boiling points Cl₂ −34°C, Br₂ 59°C, I₂ 184°C illustrate trend/data interpolation; e.g. bromine is liquid at room temperature because its boiling point is above room temperature, while chlorine is gas. Chlorine test: damp blue litmus turns red then is bleached. Halogens react with metals to make halides: 2Na+Cl₂→2NaCl; hydrogen reacts to make hydrogen halides, H₂+Cl₂→2HCl and H₂+Br₂→2HBr, which dissolve in water to make acidic solutions. Predict unfamiliar examples from the trends.
A more reactive halogen displaces a less reactive halogen from a solution of its halide; e.g. chlorine added to potassium bromide forms potassium chloride and bromine. The trend predicts that bromine cannot displace chloride and that astatine is less reactive than iodine.
Halogens have seven electrons in their outer shell. Down Group 7, additional occupied shells increase distance and shielding; the nucleus attracts an incoming electron less strongly, so gaining one electron becomes less easy and reactivity falls.
Noble gases have full outer shells and are chemically inert compared with elements that readily gain, lose or share electrons. Their uses match their properties: helium's low density for balloons, argon's inertness for shielding hot metal, and non-flammable gases where fire risk matters. Boiling points increase down Group 0 as atoms get larger and intermolecular attractions strengthen; for example He about −269°C, Ne −246°C, Ar −186°C, Kr −153°C. Read a supplied data table and extrapolate cautiously to an unknown noble gas.
What to do in an assessment
State the observed trend, compare both named substances and use outer electron arrangement or intermolecular attractions to explain the trend.
For displacement, compare halogen positions in the trend and predict whether a reaction occurs and which halogen is released.
In plain language
A group's position points to outer-electron structure. That structure helps explain reactivity trends and the physical properties of its elements.
Self-check
Will bromine displace chloride ions from potassium chloride?
Show the answer
No. Chlorine is more reactive than bromine, so bromine does not displace chlorine from chloride.
Explain reaction rates and energy changesCore
Scope: Core · Topic 7, points 7.1–7.16; core practical 7.1
Open revision guidance
What to know
Particles react when they collide with sufficient energy. Higher temperature raises particle kinetic energy; higher concentration or gas pressure increases collision frequency; greater solid surface area exposes more particles. Each can increase rate.
Choose a rate method to fit the reaction: use a gas syringe for gas production, balance/mass for gas escaping, colour/disappearing cross for visible precipitate, or sample-and-titrate for reactant concentration. Core practical 7.1 part (a): add marble chips to HCl, measure collected gas volume at regular times, plot gas volume (vertical) against time (horizontal) and compare initial gradients while varying acid concentration, temperature or marble-chip surface area. Part (b): warm sodium thiosulfate, add dilute HCl over a marked cross, time until sulfur precipitate obscures it; repeat at different temperatures and plot time (vertical) against temperature (horizontal). In both parts control all other variables, use repeats and a consistent endpoint, and wear eye protection; thiosulfate reaction makes sulfur dioxide, so ensure ventilation and do not inhale fumes.
Higher temperature makes particles move faster and raises collision energy; increased concentration/pressure raises collision frequency; a greater solid surface-area-to-volume ratio gives more collision sites. More frequent collisions and/or a greater proportion above activation energy increase rate. For a cube of side l, surface area = 6l² and volume = l³, so SA:V = 6/l; halving l doubles area per volume. For point 7.5, interpret mass-, gas-volume- and concentration-against-time graphs: e.g. flask contents losing 4.0 g in 40 s has a mass-loss rate 0.10 g s⁻¹; a reactant falling from 0.80 to 0.60 mol dm⁻³ over 50 s changes at 0.004 mol dm⁻³ s⁻¹ (use gradient magnitude for rate); a 60 cm³ gas increase over 40 s is 1.5 cm³ s⁻¹. Use two points or a tangent/gradient for a rate, not a single coordinate. Steeper tangent means faster reaction, initial tangent gives initial rate, plateau means reaction ended. For gas-volume-versus-time, volume is vertical (y) and time is horizontal (x).
A catalyst speeds reaction without changing products and is chemically unchanged and remains the same mass at the end of the reaction (when recovered). It provides an alternative route with lower activation energy, so a greater fraction of collisions are successful at the same temperature. Enzymes are biological catalysts, including those used by yeast in alcoholic-drink production.
Heat changes occur when salts dissolve, and in neutralisation, displacement and precipitation; solution temperature change evidences heat transfer. For examples, dissolve ammonium chloride in water (temperature falls), neutralise hydrochloric acid with sodium hydroxide (temperature rises), displace copper from copper sulfate with magnesium (temperature rises), and mix barium chloride with sodium sulfate to precipitate barium sulfate (record the observed change). For each, measure the same solution volume's initial temperature, add a measured reactant amount, stir, record the highest/lowest final temperature and calculate ΔT; use an insulated cup, repeat and keep amounts/starting conditions controlled. Exothermic changes release heat; endothermic changes absorb heat. Breaking bonds requires energy and making bonds releases energy. Activation energy is the minimum energy needed for a reaction to occur.
Reaction-profile diagrams plot energy against reaction progress: label reactants/products, activation-energy arrow from reactant energy level up to peak, and signed ΔH arrow from reactant level to product level. Products lower than reactants give negative ΔH/exothermic; products higher give positive ΔH/endothermic. Compare each change with a zero baseline. To investigate heat changes for a salt dissolving, neutralisation, displacement and precipitation, use an insulated polystyrene cup, record starting solution temperature, add a measured amount of salt/reactants, stir and record highest/lowest temperature and ΔT; repeat consistently and compare sign/magnitude, noting heat lost to cup/surroundings and thermometer resolution.
What to do in an assessment
Calculate mean rate as measured change ÷ time; find a graph gradient or tangent with units. For a single plotted coordinate, report the variable value at that time; a rate needs at least a time interval or graph gradient/tangent.
For point 7.5, read mass-, gas-volume- or concentration-versus-time graphs and interpret slope in the corresponding units; gas volume is plotted vertically against time horizontally.
In a practical plan repeats and controls, choose a suitable range, record times/volumes precisely and evaluate uncertainty or a subjective colour end point.
Use relative reactant/product energy levels to classify the reaction and label activation energy.
In plain language
Rates depend on successful collisions, while energy profiles show the energy barrier and overall heat transfer.
Self-check
What does a steeper gas-volume-versus-time graph show?
Show the answer
A greater rate of gas production at that time.
Evaluate fuels, atmosphere and Earth chemistryCore
Scope: Core · Topic 8, points 8.1–8.26
Open revision guidance
What to know
A hydrocarbon contains only carbon and hydrogen. Crude oil is a finite, complex mixture with chain or ring molecules and is an important fuel and petrochemical feedstock. Fractional distillation separates it by boiling-point range: vaporise crude oil, then vapours condense at different heights in a temperature-gradient column. Fractions are mixtures, not pure compounds.
Fractions from top to bottom include gases (domestic heating/cooking), petrol (cars), kerosene (aircraft), diesel (some cars/trains), fuel oil (ships/power stations) and bitumen (roads/roofing). In a supplied comparison table, trace the trend from gases to residue: carbon/hydrogen chain and molecule size increase; boiling range and viscosity increase; ignition becomes harder. Link viscosity to flow/pumping and boiling range to column collection/use, and cite the table values. Alkanes form a homologous series: same general formula CₙH₂ₙ₊₂, adjacent members differ by CH₂, boiling point rises gradually as chains lengthen, and similar chemistry follows from the shared single-bond structure.
Complete combustion in oxygen releases energy and makes CO₂ + H₂O; e.g. CH₄ + 2O₂ → CO₂ + 2H₂O. Limited oxygen gives incomplete combustion, producing carbon monoxide and/or carbon soot. CO is toxic because it binds haemoglobin and reduces oxygen transport; soot harms health and soils appliances. Sulfur impurities form SO₂, which dissolves/oxidises in rainwater to acidify soils/lakes and damage organisms/buildings. At engine temperatures N₂ and O₂ react to form pollutant nitrogen oxides.
Compare hydrogen with petrol as a car/bus fuel using energy per use, exhaust products, storage/transport, safety, cost and production source: hydrogen use produces water, but is not automatically low-carbon if made using fossil fuels; compressed storage and distribution need infrastructure. Petrol produces CO₂ and pollutants on combustion but has an established liquid-fuel network. In a data question compare emissions per distance/energy and identify where each figure is measured; include hydrogen manufacture and storage, not only tailpipe. Petrol, kerosene and diesel are non-renewable crude-oil fuels; methane is non-renewable natural gas.
Cracking breaks long-chain saturated alkanes into shorter, more useful molecules, including alkenes: large alkane → smaller alkane + alkene. It meets demand for short-chain fuels and alkene feedstocks because crude oil contains more long-chain fractions than desired. Set up a hard-glass test tube with paraffin-soaked mineral wool at one end and heated catalyst/porous ceramic chips downstream; clamp, heat catalyst first, then vaporise paraffin over it and collect gas over water. Test a sample with bromine water (decolourises if alkene formed). Exclude oxygen, use small quantities, eye protection and ventilation, keep flammable vapour away from ignition except at controlled heating, remove delivery tube before cooling to avoid suck-back, and stop before the tube runs dry.
Early atmosphere formed from volcanic gases and is thought to have contained little/no oxygen, much carbon dioxide, water vapour and small amounts of other gases. As Earth cooled, water vapour condensed to oceans and CO₂ dissolved; primitive plants used CO₂ in photosynthesis and released O₂, gradually increasing oxygen. Oxygen is tested with a glowing splint that relights.
Carbon dioxide, methane and water vapour absorb outgoing infrared radiation and re-emit energy, keeping Earth warmer (greenhouse effect). Today’s dry atmosphere is about 78% nitrogen, 21% oxygen and roughly 1% other gases, including argon and variable trace CO₂. Evaluate evidence for human climate influence by comparing CO₂ records, fossil-fuel use and temperature trends; correlation supports but alone does not prove causation. Consider measurement location, historical accuracy, timescale and uncertainty. Increased CO₂ from fossil fuels and methane from livestock may alter climate; assess mitigation by scale, risk, cost and environmental consequences.
What to do in an assessment
Use balanced equations for combustion and distinguish CO toxicity from soot/particulates; identify pollutant and source.
When comparing fractions, name the direction of change for molecule size, carbon/hydrogen number, boiling range, ignition ease and viscosity, then cite table evidence.
When interpreting climate or fuel data, quote evidence, identify a limitation and make a conclusion proportionate to the data. For a fuel choice, include full lifecycle manufacture/distribution and use-specific requirements.
In plain language
Fuel choices affect both energy supply and the air. Earth's atmosphere changed naturally, while evidence can also be used to assess human influence.
Self-check
Why can incomplete combustion of a hydrocarbon be dangerous?
Show the answer
It can produce carbon monoxide, a toxic gas that reduces the blood's ability to carry oxygen.
Identify ions and explain organic materialsCore
Scope: Core · Topic 9, points 9.1C–9.39C; core practicals 9.6C and 9.28C
Open revision guidance
What to know
A qualitative ion test identifies an ion by a result that separates it from plausible alternatives. Flame-test colours: Li⁺ red, Na⁺ yellow, K⁺ lilac, Ca²⁺ orange-red, Cu²⁺ blue-green. Clean a wire loop between samples. Add sodium hydroxide dropwise, then in excess: Al³⁺ white precipitate dissolves in excess; Ca²⁺ white precipitate remains; Cu²⁺ blue; Fe²⁺ green turning brown; Fe³⁺ brown. Warm ammonium salt with NaOH; ammonia turns damp red litmus blue. Use a clean sample and appropriate small volume so precipitate formation/dissolution can be distinguished.
For anions: add dilute acid to carbonate and test evolved CO₂ with limewater (milky); acidify with dilute HCl then add barium chloride for sulfate (white precipitate); acidify with dilute nitric acid then add silver nitrate for halides: chloride white, bromide cream, iodide yellow. Worked unknown: an orange-red flame and persistent white NaOH precipitate support Ca²⁺; a white precipitate after acidifying a fresh portion with HCl and adding BaCl₂ supports SO₄²⁻, so the salt is calcium sulfate. Confirm by combining independent positive results and checking competing ions; acidify sulfate/halide portions in the specified order to prevent carbonate interference. Core practical 9.6C: test unknown salts, record reagent/method and positive/negative result for each, infer both ions; use clean apparatus and goggles.
Flame photometry is instrumental and can be more sensitive/accurate and faster for metal ions than visual flame colour, particularly at low concentration; it requires calibration/standards and suitable instrument. Example calibration data: standard concentrations 0.10, 0.20 mol dm⁻³ give readings 20, 40 units; an unknown reading of 30 units interpolates halfway to 0.15 mol dm⁻³. Separately identify an unknown metal ion by comparing its measured flame-photometer reading with supplied reference readings for known ions collected at the same instrument settings. Worked interpretation: if the unknown reading matches, or is nearest to, the K⁺ reference row rather than the Li⁺ or Na⁺ rows, identify K⁺ as the likely match; cite the compared table entries and note overlap/resolution if relevant. Plot reading (vertical) against concentration (horizontal) for a calibration curve; for identification, compare unknown readings with reference-ion readings, not the calibration curve alone. No instrument-mechanism knowledge is needed.
Connectivity notation (not a displayed all-bond drawing): alkanes are saturated and have only C–C single bonds: methane H–C(H)(H)–H (CH₄); ethane H–C(H)(H)–C(H)(H)–H; propane H–C(H)(H)–C(H)(H)–C(H)(H)–H; butane H–C(H)(H)–C(H)(H)–C(H)(H)–C(H)(H)–H (general formula CₙH₂ₙ₊₂). In this notation, each parenthesised H is attached to the adjacent C and each dash joins the neighbouring atoms; the alkane-displayed-formulas visual gives the actual all-atom/all-bond drawings. Alkenes are unsaturated with C=C: ethene H–C(H)=C(H)–H; propene H–C(H)=C(H)–C(H)(H)–H; but-1-ene H–C(H)=C(H)–C(H)(H)–C(H)(H)–H; but-2-ene H–C(H)(H)–C(H)=C(H)–C(H)(H)–H (general formula CₙH₂ₙ). In this connectivity notation each carbon has four covalent bonds; use the alkene-displayed-formulas visual for actual all-atom/all-bond drawings and the differing butene double-bond positions. Addition: CH₂=CH₂+Br₂→CH₂Br–CH₂Br (1,2-dibromoethane); apply to propene by opening its double bond and adding one Br to each carbon. Bromine water changes orange to colourless with alkene. Complete combustion is oxidation: e.g. CH₄+2O₂→CO₂+2H₂O; C₂H₄+3O₂→2CO₂+2H₂O.
A polymer is a substance with high average relative molecular mass made of repeat units. Addition polymerisation joins alkene C=C bonds: poly(ethene) [–CH₂–CH₂–]ₙ; poly(propene) [–CH₂–CH(CH₃)–]ₙ; PVC [–CH₂–CHCl–]ₙ from chloroethene; PTFE [–CF₂–CF₂–]ₙ from tetrafluoroethene. To infer monomer from a repeat unit, identify adjacent backbone carbons, restore C=C and remove continuation bonds. Relate flexible, tough PE/PP to containers/ropes; PVC's durable electrical insulation to cable coating; PTFE's low-friction chemical resistance to non-stick/chemical tubing. Compare a supplied property table with the use requirement rather than relying on label alone.
Polyester forms by condensation of a dicarboxylic acid and a diol: HOOC–C₆H₄–COOH + HO–CH₂–CH₂–OH forms repeating [–CO–C₆H₄–CO–O–CH₂–CH₂–O–]ₙ ester-linked units; each –COOH/–OH link forms with one H₂O removed. Mark the –COO– linkage and lost H/OH atoms on a structural drawing. Polymer issues: finite fossil feedstock, persistent landfill, combustion gases, and sorting needed before remelting. Evaluate recycling using feedstock availability, sorting/processing economics, quality loss and environmental effects; DNA has four nucleotide monomers, starch is sugar-based, proteins amino-acid polymers.
Connectivity notation (not displayed all-bond drawings): methanol H–C(H)(H)–O–H, ethanol H–C(H)(H)–C(H)(H)–O–H, propan-1-ol H–C(H)(H)–C(H)(H)–C(H)(H)–O–H and butan-1-ol H–C(H)(H)–C(H)(H)–C(H)(H)–C(H)(H)–O–H; –OH is the alcohol functional group. Use the alcohol-displayed-formulas visual for actual drawings with every atom and bond, including C–O and O–H. Dehydration removes H₂O to form an alkene. Carboxylic-acid connectivity notation: methanoic H–C(=O)–O–H, ethanoic H–C(H)(H)–C(=O)–O–H, propanoic H–C(H)(H)–C(H)(H)–C(=O)–O–H and butanoic H–C(H)(H)–C(H)(H)–C(H)(H)–C(=O)–O–H; –COOH gives acidic aqueous solutions. Use the carboxylic-acid-displayed-formulas visual for every atom and C–H, C–C, C=O, C–O and O–H bond. Ethanol oxidises to ethanoic acid. Members share functional group/similar reactions; successive members differ by CH₂.
Core practical 9.28C: burn ethanol, propanol, butanol and pentanol beneath a calorimeter containing a known mass of water; record initial and highest final water temperature, calculate ΔT, and weigh each burner before/after to find mass of fuel burned. Keep water mass, burner distance, wick and draught conditions consistent; repeat, stir water, use heat shield and eye protection, keep flammable alcohols away from ignition and extinguish safely. Compare temperature rise for the same water mass and fuel burned; explain that heat loss and incomplete combustion affect results.
Fermentation uses yeast enzymes to convert glucose/carbohydrate in aqueous solution to ethanol and CO₂ in warm, oxygen-limited conditions. Fermentation mixture contains dilute ethanol; fractional distillation concentrates it because ethanol has a lower boiling point than water. In a fractionating-column set-up use a flask, column, thermometer and water-cooled condenser; warm safely without a naked flame near ethanol, collect ethanol-rich distillate as temperature approaches its boiling point and leave most water in flask.
Nanoparticles are 1–100 nm; atom diameter is about 0.1 nm=10⁻¹⁰ m, so nanoparticles are roughly 10–1000 atom diameters across (molecules vary in size). Write 10 nm=1×10⁻⁸ m and compare powers of ten. For a cube side L, SA=6L² and V=L³, so SA:V=6/L: at L=100 nm it is 0.06 nm⁻¹; at 10 nm it is 0.6 nm⁻¹, ten times greater. Smaller particles may show different reactivity/properties; sunscreen nanoparticles can absorb/reflect UV while appearing transparent. Assess possible exposure to tissues/air/water and uncertainty, not a universal harm claim. Approximate typical densities (formulation varies): soda-lime glass 2.5, fired clay ceramic 2.2, polystyrene 1.05, glass-fibre polyester composite 1.8, aluminium 2.7, copper 8.9 g cm⁻³. For a light packaging insert not required to carry a structural load, polystyrene is the lowest-density listed and insulating; for a rigid transparent window, glass is transparent but a brittle/impact trade-off matters. In an exam, compare all supplied strength, density, flexibility, conductivity, durability and cost figures, then cite values and trade-offs for the specific use rather than extrapolating density alone.
Size comparison for 9.35C: a small molecule such as H₂O is roughly 0.3 nm across, compared with an atom diameter of about 0.1 nm and nanoparticles 1–100 nm. Thus even a 1 nm nanoparticle is roughly ten atom diameters and several small-molecule widths across; a 10 nm particle is about 100 atom diameters or 30 water-molecule widths. Molecules vary in size, so state this as a qualified example, not a universal molecular size.
What to do in an assessment
For an unknown salt, test fresh portions for cation and anion, record reagent/order/positive and negative observations, then use the combined evidence to name both ions and salt.
Use the displayed-formula diagrams to trace all atoms and bonds; the inline H–C(H)… expressions are connectivity notation, not displayed drawings. For polymers show repeat-unit brackets, continuation bonds and monomer inversion.
For a material choice, compare the supplied table against the required function, cite quantitative values, weigh trade-offs and state a supported selection.
In plain language
Separate Chemistry extends ion analysis and organic chemistry. Tests identify particles; structures and functional groups help predict how molecules react and what materials can do.
Self-check
What observation with bromine water supports the presence of an alkene?
Show the answer
Orange bromine water is decolourised as bromine adds across the carbon–carbon double bond.
View diagrams
Alkanes: displayed formulae
Methane · CH₄
Every carbon and hydrogen atom and every C-C and C-H bond is shown; each carbon has four covalent bonds.
Ethane · C₂H₆
Every carbon and hydrogen atom and every C-C and C-H bond is shown; each carbon has four covalent bonds.
Propane · C₃H₈
Every carbon and hydrogen atom and every C-C and C-H bond is shown; each carbon has four covalent bonds.
Butane · C₄H₁₀
Every carbon and hydrogen atom and every C-C and C-H bond is shown; each carbon has four covalent bonds.
Every carbon and hydrogen atom is labelled, and every C-H and C-C covalent bond is drawn.
Alkenes: displayed formulae
Ethene · C₂H₄
Every atom and bond is shown; the double bond is between C1 and C2.
Propene · C₃H₆
Every atom and bond is shown; the double bond is between C1 and C2.
But-1-ene · C₄H₈
Every atom and bond is shown; the double bond is between C1 and C2.
But-2-ene · C₄H₈
Every atom and bond is shown; the double bond is between C2 and C3.
Every atom and bond is drawn. The double-bond position is shown explicitly, so but-1-ene and but-2-ene are different structures.
Alcohols: displayed formulae
Methanol · CH₄O
Every atom is shown: the -OH group has explicit C-O and O-H bonds, and all hydrogen atoms are included.
Ethanol · C₂H₆O
Every atom is shown: the -OH group has explicit C-O and O-H bonds, and all hydrogen atoms are included.
Propan-1-ol · C₃H₈O
Every atom is shown: the -OH group has explicit C-O and O-H bonds, and all hydrogen atoms are included.
Butan-1-ol · C₄H₁₀O
Every atom is shown: the -OH group has explicit C-O and O-H bonds, and all hydrogen atoms are included.
The C-O and O-H bonds in each hydroxyl group are explicit; all carbon-bound and oxygen-bound hydrogen atoms are drawn.
Carboxylic acids: displayed formulae
Methanoic acid · CH₂O₂
Every atom is shown: the carboxyl carbon has a C=O double bond and a C-O-H group, with every hydrogen included.
Ethanoic acid · C₂H₄O₂
Every atom is shown: the carboxyl carbon has a C=O double bond and a C-O-H group, with every hydrogen included.
Propanoic acid · C₃H₆O₂
Every atom is shown: the carboxyl carbon has a C=O double bond and a C-O-H group, with every hydrogen included.
Butanoic acid · C₄H₈O₂
Every atom is shown: the carboxyl carbon has a C=O double bond and a C-O-H group, with every hydrogen included.
Each structure shows the carboxyl group with both C=O and C-O-H bonds, as well as every C-C and C-H bond.
In dot-and-cross diagrams, a dot is an electron from the atom named in the explanation and a cross is from the other atom. The symbols show electron origin, not a charge. Shared pairs belong between bonded atoms; lone pairs stay on one atom.
Use these all-atom/all-bond displayed structures to count bonds around each carbon and compare chain length, double-bond position, –OH and –COOH groups. Every hydrogen atom and its covalent bonds are drawn.
Plan practical work and reason with dataCore
Scope: Core · §2.4 and Appendix 1, Appendix 2, Appendix 4, Appendix 5; Working Scientifically; all eight core practicals
Open revision guidance
What to know
The qualification requires all eight mandatory core practicals, a student record of the work undertaken and the knowledge, skills and understanding derived from it, and a centre Practical Science Statement confirming completion. The statement is submitted to Pearson by 15 April in the examination year. Written papers also assess knowledge and understanding of practical methods.
The eight activities investigate ink separation (2.11); pH as calcium oxide/hydroxide is added to hydrochloric acid (3.6); copper sulfate crystal preparation (3.17); copper sulfate electrolysis (3.31); acid–alkali titration (5.9C); rate changes using gas volume or colour change (7.1); unknown salt ions (9.6C); and alcohol combustion temperature rise (9.28C). These are a minimum, not a cap on practical work.
Practical 2.11 uses two set-ups: ink paper chromatography with baseline above solvent and a covered chamber; simple distillation to recover pure water from ink using a heated flask, condenser with cooling water and receiver. Practical 3.6 adds weighed portions of calcium oxide/hydroxide to fixed dilute HCl, recording pH after each addition by meter or universal indicator paper/glass rod. Practical 3.17 warms dilute sulfuric acid in a water bath, adds excess copper oxide, filters, evaporates solution and dries hydrated blue crystals.
Practical 3.31 varies current during CuSO₄ electrolysis with copper electrodes and measures electrode mass change; separately observes products with graphite electrodes. Practical 5.9C uses burette, pipette, indicator, white tile and swirling to reach the endpoint, read/record titres and identify concordant results. Foundation learners' required practical outcome is the accurate method, endpoint and results; mole-ratio calculations for unknown concentration or required volume are Higher-only 5.10C. Practical 7.1 follows marble-chip/HCl gas volume against time and sodium-thiosulfate/HCl disappearing-cross time at varied conditions; plot/process rates. Practical 9.6C tests unknown salts using specified flame, precipitation and gas procedures. Practical 9.28C measures temperature rise of a known water mass as four alcohols burn. The specification appendix's procedure details must be followed; risk controls and eye protection apply.
AT1 accurate mass, time, temperature and liquid/gas volume; AT2 safe heating including Bunsen burner and water bath/electric heater; AT3 conducting/monitoring reactions and measuring pH; AT4 evaporation, filtration, crystallisation, chromatography and distillation; AT5 observations, temperature changes and rate; AT6 safe handling/controlled mixing of gases, liquids, solids/reagents; AT7 electrochemical cells; AT8 gas/flame/precipitation identification tests and determining concentrations of strong acids and strong alkalis. For AT8 practise accurate titration apparatus/method, endpoint and recording results; its mole-ratio concentration calculations are the Higher-only 5.10C supplement, not Foundation content. pH measurement belongs to AT3, including Core Practical 3.6.
Chemistry maths 1a–1d: write 0.25 as a decimal (1a); 6.02×10²³ in standard form (1b); express a 1:3 mole ratio as a fraction/percentage (1c); estimate whether 2.0 g Mg could make about 3 g MgO before calculating (1d).
Maths 2a, 2b, 2c, 2h: report 2.436 g to 3 significant figures as 2.44 g (2a); average concordant titres 24.8, 24.9, 24.9 cm³ to 24.9 cm³ (2b); construct a frequency table/bar chart for reaction times and a touching-bin histogram for continuous temperature data, then interpret a common interval (2c); compare nanoparticle diameter 10⁻⁸ m with a 10⁻¹⁰ m atom as two orders of magnitude larger (2h).
Maths 3a–3c: use =, <, <<, >>, >, ∝ and ~ correctly (e.g. at fixed conditions gas volume ∝ amount); rearrange c=m/V to V=m/c (3b); substitute m=5.0 g and V=0.25 dm³ to obtain c=20 g dm⁻³, retaining units (3c).
Maths 4a–4e: translate a graph coordinate (40 s, 60 cm³) into the numerical statement that gas volume is 60 cm³ at 40 s; one point alone is not a rate (4a); interpret y=mx+c as a linear relation (4b); plot gas volume on the vertical axis against time on the horizontal axis, with labelled, uniform scales and best-fit line (4c); a calibration line through (0,0) and (0.20 mol dm⁻³, 40 units) has gradient 200 units dm³ mol⁻¹ and intercept 0 (4d); draw a tangent on a curved rate graph and divide rise by run with units for instantaneous rate (4e).
Maths 5b–5c: convert a 3-D ionic lattice or molecule into a labelled 2-D representation and explain what depth information is lost (5b); triangle base 4 cm and perpendicular height 3 cm has area ½×4×3=6 cm², a 4×3 rectangle is 12 cm², and a 2 cm cube has SA=24 cm² and V=8 cm³ (5c). Chemistry does not list 2d–g, 3d, 4f or 5a as assessed chemistry skills.
Working Scientifically 1a–1f: explain the atom-model sequence and how new evidence changed a theory (1a); choose a particle/model/graph representation to predict a reaction (1b); state a model's limits and ethical issues of a process (1c); evaluate a fuel/material's personal, social, economic and environmental effects using evidence (1d); assess practical and societal risks by probability and consequence, not just perceived danger (1e); explain why results are peer-reviewed and communicate a finding to a class/public audience (1f).
Working Scientifically 2a–2g: form a testable hypothesis from collision theory (2a); plan a rate experiment with one changed factor and measured outcome (2b); select gas syringe, balance or stopwatch to suit the endpoint (2c); use measuring apparatus correctly and follow hazards/control measures (2d); take samples across a heterogeneous solid or repeated flame tests to avoid a biased sample (2e); record mass, time, temperature and volume with resolution and units (2f); evaluate a method, suggest a specific improvement and a follow-up test (2g).
Working Scientifically 3a–3h: present observations in a table/graph (3a); translate a word equation or graph into a symbol/numeric form (3b); calculate means, ratios, gradients and rate statistics (3c); show spread/range and measurement uncertainty (3d); infer a trend from numerical, verbal, diagrammatic or symbolic evidence (3e); connect a conclusion to a hypothesis and data (3f); distinguish accuracy, precision, repeatability and reproducibility and identify random/systematic errors without cherry-picking (3g); report rationale, method, findings and evidence-based conclusion in a labelled diagram, table, graph and prose (3h).
Working Scientifically 4a–4c: use precise terms such as precipitate, electrolyte, concentration and exothermic (4a); recognise how mass, time, temperature and volume are measured (4b); use SI units and IUPAC names/formulae (4c). Appendix 1, not Working Scientifically, covers mathematical powers of ten (Maths 1b), significant figures (2a) and substitution/units (3c); for example 250 cm³=0.250 dm³. Issue 4 has no Working Scientifically clauses 4d–4f.
Appendix 4 technique-to-practical map: AT1 accurate mass, time, temperature and liquid/gas volume (5.9C burette readings; 9.28C water temperature/fuel mass); AT2 safe Bunsen and water-bath/electric-heater use (3.17 crystals; 7.1 controlled warming); AT3 conduct/monitor a reaction and measure pH (3.6); AT4 evaporate, filter, crystallise, chromatograph and distil (2.11 and 3.17); AT5 observe colour/gas/solid/temperature and measure rate (3.31, 7.1); AT6 safely handle and control mixing of gases, liquids, solids/reagents (all practicals, especially acids and alcohols); AT7 set up/use electrochemical cells with electrodes, leads and DC supply (3.31); AT8 qualitative gas, flame and precipitation tests (9.6C) and determine strong-acid/strong-alkali concentration by titration (5.9C). pH measurement is AT3, not AT8. Foundation candidates practise the common titration method/end-point/data; mole-ratio calculations for unknown concentration or volume are Higher-only 5.10C.
Practical methods and records: 2.11 mark pencil baseline above solvent, spot/dry ink, develop and mark solvent front; distil ink with anti-bumping granules, condenser cooling water entering low and leaving high, and an unsealed receiver. 3.6 add weighed powdered CaO/Ca(OH)₂ portions to fixed dilute HCl, stir, measure/record pH after each and graph pH against cumulative solid mass. 3.17 warm dilute H₂SO₄ in a water bath, add CuO until excess, filter, evaporate filtrate to near saturation, cool to crystallise, filter and dry blue hydrated crystals. 3.31 use separate inert/copper electrodes in CuSO₄, record products and mass changes, vary/record current and switch off before handling. 5.9C use pipette, burette, suitable indicator and concordant titres. 7.1 measure marble/HCl gas at intervals and thiosulfate/HCl disappearing-cross time. 9.6C identify ions by specified flame, precipitate and gas tests. 9.28C burn four alcohols beneath calorimeter with known water mass, record temperature rise and burner mass change. Keep a labelled apparatus drawing, readings, observations, calculations, safety controls and conclusion in each learner practical record; every one of the eight is mandatory, additional practicals are encouraged, and written exams assess method knowledge.
Appendix 2 command-word meanings: Add/Label adds a label/unit to supplied material; Assess weighs relevant factors and judges importance; Calculate obtains a number with working and unit; Comment on synthesises variables into a judgement; Compare gives similarities/differences for all named items without requiring a conclusion; Compare and contrast gives both similarities and differences, also without requiring a conclusion; Complete fills a table/diagram; Deduce reaches a conclusion from supplied information; Describe gives a developed account without a required reason; Determine reaches an answer quantitatively from stimulus/working; Devise plans a procedure; Discuss identifies the issue and explores aspects of it through reasoning/argument; Draw produces a diagram with ruler/freehand as appropriate; Estimate finds an approximate value; Evaluate reviews information, strengths/weaknesses or alternatives and gives an evidence-supported judgement in context; Explain justifies a point; Give/State/Name recalls information; Give a reason/reasons gives why a statement is true; Identify selects key information from a stimulus; Justify supplies evidence for a claim/answer; Measure determines dimensions/angle with an instrument; Plot accurately marks supplied points, chooses labelled suitable scales and draws best-fit line; Predict states an expected result; Show that verifies a given statement; Sketch makes an unscaled freehand graph/diagram showing important features; State and explain makes a point and links reasoning; State what is meant by defines a term where different descriptions are possible; Write gives an equation. ‘Suggest a…’ asks for a proposed explanation or description.
What to do in an assessment
For each investigation, state independent/dependent/control variables, method, apparatus and range; choose a suitable resolution, repeat measurements where appropriate, record units, keep controls fixed and plan safe handling. Draw labelled apparatus diagrams when useful.
Process data using means, ratios/percentages, suitable tables/graphs and significant figures; plot labelled axes with sensible scales and best-fit line, use slope or tangent gradient with units, and interpret limits. Separate accuracy from precision, repeatability from reproducibility, and random from systematic error; propose a feasible change and explain its likely benefit.
Read the command word before composing an answer. For compare, address both/all named items; for compare and contrast, include a similarity and a difference; for calculate/show that/determine, show quantitative working and required unit; for plot, label axes, scale and best-fit line; for evaluate/assess/comment on, use evidence and a supported judgement; for explain/state and explain, connect claim and reason; for write, supply the required equation.
In plain language
Good practical science connects a safe method to reliable measurements, careful calculations and conclusions that do not claim more than the evidence shows.
Self-check
What is the difference between repeatability and reproducibility?
Show the answer
Repeatability is consistency under the same method, equipment and operator; reproducibility is consistency when relevant conditions such as operator or equipment change.
Higher: write balanced ionic equationsHigher
Scope: Higher · Paper tier: Higher tier · Higher-only point 0.4
Open revision guidance
What to know
An ionic equation keeps only particles that change; spectator ions that remain unchanged on both sides are cancelled. Conserve atoms and total charge. Example: Ag⁺(aq) + NO₃⁻(aq) + Na⁺(aq) + Cl⁻(aq) → AgCl(s) + Na⁺(aq) + NO₃⁻(aq); cancel Na⁺ and NO₃⁻ to obtain Ag⁺(aq) + Cl⁻(aq) → AgCl(s).
What to do in an assessment
Write complete balanced equation with states, split aqueous ionic substances into ions, cancel identical spectator ions, then check atom and charge totals.
In plain language
A net ionic equation shows just the particles that take part in the chemical change.
Self-check
Write the net ionic equation for neutralising aqueous hydrochloric acid with aqueous sodium hydroxide.
Show the answer
H⁺(aq) + OH⁻(aq) → H₂O(l). Na⁺ and Cl⁻ are spectator ions.
Higher: calculate isotope masses, moles and reacting amountsHigher
Scope: Higher · Paper tier: Higher tier · Higher-only points 1.12 and 1.50–1.53
Open revision guidance
What to know
Relative atomic mass from isotopes is the weighted mean: sum(isotope mass × percentage abundance) ÷ 100. Example: 75% mass 35 and 25% mass 37 gives Ar=(75×35+25×37)/100=35.5. One mole is 6.02×10²³ atoms, molecules, formula units or ions, and its mass in grams is numerically the relative particle mass.
Use n=m/Mr, m=nMr, N=nL and n=N/L, with L=6.02×10²³ mol⁻¹. These equations convert mass↔moles, moles↔particles and mass↔particles; use standard form carefully and keep units. A balanced equation gives mole ratios. The reactant not in excess limits product amount: calculate available moles and compare n/coefficient for each reactant. Deduce stoichiometric ratios from measured masses by converting each to moles and simplifying.
Example: N₂+3H₂→2NH₃. If 2 mol N₂ and 3 mol H₂ are mixed, H₂ supports 1 mol N₂ reaction and is limiting; it forms 2 mol NH₃ while 1 mol N₂ remains. Convert to mass with Mr only after applying the coefficient ratio.
What to do in an assessment
Show isotope weighting, mole conversion, equation balancing and coefficient ratio as separate steps; retain unrounded working until the final answer.
For a limiting-reactant task, divide each reactant amount by its equation coefficient; the smaller value controls reaction extent. Calculate product amount from that reactant.
In plain language
Higher quantitative chemistry links particle counts, isotope abundances and reacting amounts using mole ratios.
Self-check
An element has isotopes of mass 10 (20%) and mass 11 (80%). Find its relative atomic mass; how many particles are in 0.50 mol?
Show the answer
Ar=(10×20+11×80)/100=10.8. N=0.50×6.02×10²³=3.01×10²³ particles.
Higher: relate concentration to pH and explain electron transferHigher
Scope: Higher · Paper tier: Higher tier · Higher-only points 3.4, 3.5, 3.7, 3.8 and 3.27–3.29
Open revision guidance
What to know
Dilute/concentrated describes amount of solute per volume; strong/weak describes degree of ionisation. A strong acid dissociates extensively in water, weak acid only partially; a concentrated solution can still be weak. Higher hydrogen-ion concentration means lower pH; higher hydroxide-ion concentration means higher pH. A tenfold rise in [H⁺] decreases pH by 1. Use a pH meter or universal indicator consistently to compare strong and weak acids or supplied acid/alkali samples; explain the result using degree of ionisation rather than confusing strength with concentration. pH measurement is AT3, not AT8.
Anode is where oxidation occurs and cathode where reduction occurs. Oxidation is electron loss; reduction is electron gain. Write half-equations with electrons: Cu²⁺+2e⁻→Cu at cathode (reduction); 2Cl⁻→Cl₂+2e⁻ at anode (oxidation). Check atoms and charge on each side.
What to do in an assessment
Distinguish acid strength from concentration in explanations. Calculate the pH direction and one-unit change for a tenfold [H⁺] change.
For electrolysis, label anode/cathode and electron direction, write half-equations and identify oxidation/reduction by electron loss/gain.
In plain language
Higher chemistry connects ion concentration to pH and uses electron transfer to explain electrode reactions.
Self-check
[H⁺] rises tenfold. What happens to pH? In 2Cl⁻→Cl₂+2e⁻, what is oxidised?
Show the answer
pH decreases by 1. Chloride ions lose electrons, so chloride is oxidised.
Higher: explain redox displacement and changing equilibriaHigher
Scope: Higher · Paper tier: Higher tier · Higher-only points 4.2, 4.8 and 4.17
Open revision guidance
What to know
Displacement is redox: in Mg+Cu²⁺→Mg²⁺+Cu, Mg loses electrons (oxidised) and Cu²⁺ gains electrons (reduced). Link electron transfer to the observed replacement and balance charge.
Bioleaching uses bacteria to extract metal compounds into solution; phytoextraction grows plants that take up metal compounds, then biomass is harvested/processed. Evaluate each against ore grade, time, land, recovery yield, energy, cost, pollution and later extraction steps using the evidence given; neither method is automatically best.
At dynamic equilibrium in a closed system, equal forward/reverse rates maintain constant concentrations. Adding a reactant or product favours the direction that consumes it; removing one favours the direction that replaces it. In N₂+3H₂⇌2NH₃, adding N₂ shifts right but removing N₂ shifts left; adding NH₃ shifts left but removing NH₃ shifts right. Increased pressure favours fewer gas molecules, while decreased pressure favours more: higher pressure favours NH₃ (four gas moles→two), whereas lower pressure favours N₂ and H₂. Higher temperature favours the endothermic reverse direction and lower temperature the exothermic forward direction. State the position shift separately from the rate of attaining equilibrium.
What to do in an assessment
Write electron gain/loss for both species in a displacement reaction and identify which is oxidised/reduced.
Compare bioextraction methods with criteria and data, explaining a supported choice.
For equilibrium, consider both adding and removing reactants/products, count gas molecules for increases and decreases in pressure, identify the endothermic direction, and predict the position shift separately from reaction rate.
In plain language
Higher responses explain changes in electron transfer, evaluate extraction methods and predict the favoured equilibrium direction.
Self-check
For N₂+3H₂⇌2NH₃, what happens if ammonia is removed or pressure is decreased? Name the oxidised species in Mg+Cu²⁺→Mg²⁺+Cu.
Show the answer
Removing ammonia shifts right to replace it; decreased pressure shifts left to the side with more gas molecules. Magnesium loses electrons and is oxidised.
Higher: apply quantitative Separate ChemistryHigher
Scope: Higher · Paper tier: Higher tier · Higher-only points 5.8C, 5.10C, 5.15C–5.18C and 5.20C–5.21C
Open revision guidance
What to know
Molar concentration c=n/V (V in dm³); mass concentration c=m/V (g dm⁻³). Convert g dm⁻³ to mol dm⁻³ by dividing by Mr and back by multiplying. Example 9.8 g dm⁻³ NaOH (Mr 40) is 0.245 mol dm⁻³.
In a titration calculation, calculate n=cV for known solution, use balanced equation ratio, then divide by unknown volume or rearrange for required volume. Example unknown concentration: 25.0 cm³ NaOH is neutralised by 20.0 cm³ of 0.100 mol dm⁻³ HCl; HCl+NaOH is 1:1, so n(HCl)=0.100×0.0200=0.00200 mol and c(NaOH)=0.00200/0.0250=0.0800 mol dm⁻³. Example required volume: 25.0 cm³ 0.200 mol dm⁻³ H₂SO₄ reacts with NaOH by H₂SO₄+2NaOH→Na₂SO₄+2H₂O; n(H₂SO₄)=0.200×0.0250=0.00500 mol, so n(NaOH)=0.0100 mol; at 0.100 mol dm⁻³, V=0.0100/0.100=0.100 dm³=100 cm³.
Choose pathways using atom economy, yield, rate, equilibrium and useful by-products together. Worked decision pattern: if question data give route A 85% atom economy/60% yield/slow operation but a useful by-product, while route B gives 50% atom economy/90% yield/fast operation but a waste product, compare the supplied criteria and required production context; no single percentage settles the decision. State what matters most and justify it with the given values. At room temperature and pressure one mole of gas occupies 24 dm³ (24 000 cm³). Use n=V/24, coefficient ratios and m=nMr. Example CaCO₃→CaO+CO₂: 10.0 g CaCO₃ (Mr100) gives 0.100 mol CO₂, so 2.40 dm³ at r.t.p. Equal gas volumes at same T/P have equal particle numbers; coefficients give gas-volume ratios, e.g. 2H₂+O₂→2H₂O(g) has a 2:1:2 volume ratio.
Rates of attaining equilibrium increase with temperature (particles move faster and more exceed activation energy), concentration (more frequent collisions), pressure for gases (more frequent collisions) and catalyst (lower activation energy); a catalyst does not change equilibrium position. For Haber, iron catalyst and selected temperature/pressure balance raw-material and energy availability/cost against acceptable equilibrium yield in acceptable time. Higher pressure favours NH₃, while a lower temperature favours its exothermic formation but slows attainment; the operating compromise avoids impractically slow production.
What to do in an assessment
Convert all titration volumes to dm³, write the balanced mole ratio, show n=cV and units, and solve separately for unknown concentration or required volume.
For mass/gas-volume questions show mass→moles→equation ratio→gas volume; use 24 dm³ mol⁻¹ and convert cm³↔dm³ correctly.
When comparing pathways or Haber conditions, distinguish rate of attainment from equilibrium position/yield and support the trade-off using all data requested.
In plain language
Higher Separate Chemistry uses mole ratios to solve concentration/gas problems and weighs industrial yield, rate and cost.
Self-check
What volume of CO₂ at r.t.p. forms from 0.25 mol CaCO₃ in CaCO₃→CaO+CO₂?
Show the answer
The ratio is 1:1, so n(CO₂)=0.25 mol and V=0.25×24=6.0 dm³.
Higher: explain halogen displacement as redoxHigher
Scope: Higher · Paper tier: Higher tier · Higher-only point 6.12
Open revision guidance
What to know
In Cl₂+2Br⁻→2Cl⁻+Br₂, chlorine gains electrons and is reduced; bromide loses electrons and is oxidised. The more reactive halogen is reduced while the halide ion is oxidised. A displacement is therefore a redox reaction involving both electron gain and loss.
What to do in an assessment
Use a balanced ionic equation, identify the species that gains/loses electrons and name the oxidised and reduced species.
In plain language
Halogen displacement transfers electrons from halide ions to a more reactive halogen.
Self-check
In Cl₂+2I⁻→2Cl⁻+I₂, which substance is oxidised?
Show the answer
I⁻ loses electrons to form I₂, so iodide is oxidised.
Higher: calculate reaction energy from bond energiesHigher
Scope: Higher · Paper tier: Higher tier · Higher-only point 7.14
Open revision guidance
What to know
Estimate reaction energy by ΔH = total energy to break bonds in reactants − total energy released when forming product bonds, in kJ mol⁻¹. Breaking bonds is endothermic; making bonds is exothermic. Example H₂+Cl₂→2HCl with H–H=436, Cl–Cl=243, H–Cl=432 kJ mol⁻¹: ΔH=(436+243)−2(432)=−185 kJ mol⁻¹; negative means exothermic.
What to do in an assessment
Draw/count every bond broken and formed from the balanced equation, multiply each bond count by its energy, then calculate Σbroken−Σformed with kJ mol⁻¹ and interpret the sign.
In plain language
Higher bond-energy calculations compare energy absorbed to break reactant bonds with energy released as product bonds form.
Self-check
Breaking bonds requires 500 kJ mol⁻¹ in total and forming product bonds releases 620 kJ mol⁻¹. Find ΔH and classify the change.
Show the answer
ΔH=500−620=−120 kJ mol⁻¹, so it is exothermic.

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Chemistry topic list

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