Syllabus map · 22 topics · 1 HL only
IB Chemistry syllabus map
IB Chemistry is organised around two ideas, Structure and Reactivity, split into 22 topics that run from atoms and the mole to reaction mechanisms. Every topic is studied at SL and HL, with extra HL content in most of them and one topic, R1.4, that is HL only. Open any topic below for what it covers, what you need to be able to do, the terms to use precisely and the mistake that most often costs marks.
How it's assessed
| Component | SL | HL |
|---|---|---|
| Paper 1 (1A + 1B, one sitting) | 1 h 30, 55 marks, 36% | 2 h, 75 marks, 36% |
| Paper 1A: multiple choice | 30 questions, 30 marks | 40 questions, 40 marks |
| Paper 1B: data-based and experimental questions | 25 marks | 35 marks |
| Paper 2: short-answer and extended-response | 1 h 30, 50 marks, 44% | 2 h 30, 90 marks, 44% |
| Internal assessment: scientific investigation | 24 marks, 20% (3,000 words max) | 24 marks, 20% (3,000 words max) |
This is the current course, first examined in May 2025 and sat by students taking exams in 2026 and 2027. Calculators are allowed on both papers, and you get a clean copy of the chemistry data booklet in each.
Tap a topic to see what it asks, what you need to be able to do, the key terms and the trap. Topic titles follow the course structure; all explanations are ours.
Structure 1 — Models of the particulate nature of matter
How matter is built from particles: atoms and their electrons, counting particles with the mole, and the behaviour of gases.
S1.1Introduction to the particulate nature of matterSL
Classifying matter as elements, compounds and mixtures, and using the particle model to explain states of matter and changes of state. Questions ask you to classify a substance, choose a separation method or explain a heating curve in terms of particle energy.
You need to be able to
- Distinguish elements, compounds and mixtures, and match a separation technique (filtration, crystallisation, distillation, paper chromatography) to the property it relies on.
- Describe the arrangement, spacing and movement of particles in solids, liquids and gases, and use it to explain why only gases compress easily and diffuse quickly.
- Explain why the temperature stays constant on each plateau of a heating curve, and why the boiling plateau is longer than the melting one.
- Name each change of state and its reverse, and say which absorb energy and which release it.
- Relate temperature in kelvin to the average kinetic energy of the particles.
Key terms
- element
- compound
- homogeneous mixture
- heterogeneous mixture
- sublimation
- deposition
- average kinetic energy
- kelvin
Common trap Saying the temperature stops rising on a plateau because no energy is going in. Energy is still being absorbed; it goes into overcoming the attractions between particles rather than making them move faster.
Practise S1.1 in Heptavo →S1.2The nuclear atomSL
The structure of the atom, isotopes, and how mass spectra give relative atomic mass. Expect to count the particles in atoms and ions and to calculate relative atomic mass from isotope data.
You need to be able to
- Work out the numbers of protons, neutrons and electrons from a nuclide symbol, including for ions.
- State the relative mass and charge of the proton, neutron and electron, and explain why the nucleus holds nearly all the mass but almost none of the volume.
- Explain why isotopes of an element react in the same way but differ in physical properties such as density.
- Calculate relative atomic mass from a mass spectrum or abundance data, and check the answer sits nearer the more abundant isotope.
- Work backwards from a relative atomic mass to the percentage abundances of two isotopes.
Key terms
- atomic number (Z)
- mass number (A)
- isotope
- nuclide
- relative atomic mass (Ar)
- mass spectrum
- percentage abundance
Common trap Taking relative atomic mass as a simple average of the isotope masses. It is weighted by abundance, so chlorine-35 and chlorine-37 give 35.5, not 36.
Practise S1.2 in Heptavo →S1.3Electron configurationsSL
How electrons are arranged in levels, sub-levels and orbitals, and the evidence for this from emission spectra. Questions ask you to write configurations, fill orbital diagrams and interpret the hydrogen spectrum.
You need to be able to
- Explain why an emission spectrum is a series of lines, and why the lines in the hydrogen spectrum converge at higher energy.
- Write full and condensed electron configurations up to Z = 36, for atoms and ions, including the chromium and copper exceptions.
- Fill orbital diagrams using the Aufbau principle, the Pauli exclusion principle and Hund's rule, and count unpaired electrons.
- Describe the shapes of s and p orbitals and state how many orbitals the s, p and d sub-levels contain.
- HL: find the ionisation energy from the convergence limit, and use the jump in successive ionisation energies to deduce the group.
Key terms
- emission spectrum
- convergence limit
- energy level
- sub-level
- orbital
- Aufbau principle
- Pauli exclusion principle
- Hund's rule
Common trap Removing 3d electrons before 4s when forming a transition-metal ion. The 4s fills first but is also emptied first: Fe is [Ar]3d⁶4s², so Fe²⁺ is [Ar]3d⁶, not [Ar]3d⁴4s².
Practise S1.3 in Heptavo →S1.4Counting particles by mass: The moleSL
The mole as chemistry's counting unit, linking mass, number of particles and concentration. Questions are calculations: amounts, empirical and molecular formulas, and solution concentrations.
You need to be able to
- Convert between mass, amount in moles and number of particles using molar mass and the Avogadro constant.
- Find an empirical formula from percentage composition, then the molecular formula from the molar mass.
- Count the moles of each ion or atom in a given amount of compound, including formulas with brackets such as Al₂(SO₄)₃.
- Calculate concentration in mol dm⁻³ and g dm⁻³, converting cm³ to dm³, and work out dilutions.
- Use Avogadro's law: equal volumes of gases at the same temperature and pressure contain equal numbers of molecules.
Key terms
- mole
- Avogadro constant
- molar mass
- relative formula mass
- empirical formula
- molecular formula
- concentration
- standard solution
Common trap Forgetting to convert cm³ to dm³ before using n = cV. 25.0 cm³ of a 0.100 mol dm⁻³ solution holds 0.00250 mol, not 2.50 mol: divide the volume by 1000 first.
Practise S1.4 in Heptavo →S1.5Ideal gasesSL
The ideal gas model, the gas laws and the ideal gas equation, and why real gases deviate from ideal behaviour. Questions ask you to calculate pressure, volume, amount or temperature, sketch gas-law graphs and explain deviations.
You need to be able to
- State the assumptions of the ideal gas model and explain gas pressure in terms of particles colliding with the walls.
- Use pV = nRT with consistent units: pressure in Pa, volume in m³, temperature in K.
- Sketch and interpret graphs of p against V, p against 1/V, V against T and p against T.
- Use the combined gas law to find a new pressure, volume or temperature for a fixed amount of gas.
- Explain why real gases deviate most from ideal behaviour at low temperature and high pressure.
Key terms
- ideal gas
- molar volume
- absolute zero
- gas constant (R)
- ideal gas equation
- real gas
- elastic collision
Common trap Putting °C or dm³ into pV = nRT. With R = 8.31 J K⁻¹ mol⁻¹, volume must be in m³ (1 dm³ = 10⁻³ m³), pressure in Pa and temperature in kelvin, or the answer is out by powers of ten.
Practise S1.5 in Heptavo →Structure 2 — Models of bonding and structure
The ionic, covalent and metallic models of bonding, and how the type of bonding sets a material's properties.
S2.1The ionic modelSL
How ions form, how to write ionic formulas and why ionic lattices have the properties they do. Questions ask you to deduce formulas, compare melting points using lattice enthalpy and explain conductivity and brittleness.
You need to be able to
- Predict an ion's charge from its group and write formulas of ionic compounds, including polyatomic ions such as NH₄⁺, OH⁻, NO₃⁻, CO₃²⁻ and SO₄²⁻.
- Describe an ionic lattice and explain why ionic compounds have high melting points.
- Explain how lattice enthalpy depends on ionic charge and ionic radius, and use it to rank melting points (MgO far above NaCl).
- Explain why ionic compounds conduct when molten or dissolved but not as solids, and why they are brittle.
Key terms
- cation
- anion
- ionic bond
- ionic lattice
- lattice enthalpy
- polyatomic ion
- electrostatic attraction
Common trap Saying a solid ionic compound does not conduct because it has no ions. The ions are there but held in fixed positions; conduction needs mobile charged particles, which the ions only become once melted or dissolved.
Practise S2.1 in Heptavo →S2.2The covalent modelSL
The largest bonding topic: Lewis formulas, molecular shapes, polarity, covalent network structures and intermolecular forces. Questions ask you to draw structures, predict shapes and bond angles, and explain differences in boiling point or solubility.
You need to be able to
- Draw Lewis formulas, including molecules with multiple bonds and coordination bonds.
- Use VSEPR to predict shapes and bond angles for up to four electron domains, explaining why each lone pair narrows the angle (about 107° in NH₃, 104.5° in H₂O).
- Decide whether a molecule is polar from its bond polarities and its symmetry.
- Rank boiling points using London dispersion forces, dipole–dipole forces and hydrogen bonding, comparing molecules of similar size.
- Compare diamond, graphite, graphene and silicon dioxide as covalent networks; HL adds sigma and pi bonds, hybridisation, resonance and formal charge.
Key terms
- Lewis formula
- electron domain
- VSEPR
- electronegativity
- bond polarity
- London dispersion forces
- hydrogen bond
- covalent network structure
Common trap Saying covalent bonds break when a molecular substance boils. Boiling water separates whole H₂O molecules by overcoming the hydrogen bonds between them; the O–H bonds inside each molecule stay intact.
Practise S2.2 in Heptavo →S2.3The metallic modelSL
Metallic bonding as a lattice of cations in a sea of delocalised electrons, and how that model explains conductivity, malleability and melting-point trends. Questions ask you to explain properties and compare bond strength between metals.
You need to be able to
- Describe metallic bonding as the electrostatic attraction between a lattice of cations and delocalised electrons.
- Explain electrical and thermal conductivity, malleability and ductility from the model.
- Explain why melting point rises from sodium to magnesium to aluminium, using cation charge, number of delocalised electrons and ionic radius.
- Explain why the melting points of group 1 metals fall down the group.
Key terms
- metallic bond
- delocalised electrons
- cation lattice
- malleable
- ductile
- electrical conductivity
Common trap Saying metals conduct because ions move. In a solid metal the cations stay in place; it is the delocalised electrons that drift through the lattice and carry the current.
Practise S2.3 in Heptavo →S2.4From models to materialsSL
Bonding treated as a continuum, and how a material's bonding explains its properties, including alloys and polymers. Questions ask you to place compounds on the bonding triangle, explain the properties of alloys and draw polymer repeating units.
You need to be able to
- Use the difference in electronegativity and the average electronegativity to place a substance on the bonding triangle.
- Predict melting point and conductivity from the structure type: giant ionic, metallic, covalent network or simple molecular.
- Explain why an alloy is usually harder than the pure metal.
- Draw the repeating unit of an addition polymer from its monomer, and work back from a polymer chain to the monomer.
- HL: identify condensation polymers and the small molecule released as they form.
Key terms
- bonding continuum
- bonding triangle
- alloy
- polymer
- monomer
- repeating unit
- addition polymer
Common trap Drawing an addition polymer's repeating unit with the C=C still in it, or without the bonds extending from each end. The double bond becomes a single bond, and the unit sits in brackets with a continuation bond on each side.
Practise S2.4 in Heptavo →Structure 3 — Classification of matter
Sorting elements with the periodic table and organic compounds by their functional groups.
S3.1The periodic table: Classification of elementsSL
How an element's position in the periodic table follows from its electron configuration, and the trends in radius, ionisation energy and electronegativity. Questions ask you to explain trends, predict group 1 and group 17 reactions and classify oxides.
You need to be able to
- Read an element's group, period and block from its electron configuration, and the reverse.
- Explain trends in atomic and ionic radius, first ionisation energy, electron affinity and electronegativity using nuclear charge, shielding and distance.
- Explain the reactivity trends in groups 1 and 17, and write equations for halogen displacement reactions.
- Explain how oxides change from basic to acidic across a period, with equations for metal oxides and non-metal oxides such as CO₂ and SO₂ reacting with water.
- HL: explain the variable oxidation states and coloured compounds of transition elements.
Key terms
- group
- period
- block
- shielding
- first ionisation energy
- electronegativity
- electron affinity
- amphoteric
Common trap Explaining the fall in ionisation energy down a group by saying there are more protons. There are, but the outer electron is further from the nucleus and better shielded, and that outweighs the extra nuclear charge.
Practise S3.1 in Heptavo →S3.2Functional groups: Classification of organic compoundsSL
How organic compounds are represented, classified and named, and how isomers arise. Questions ask you to name compounds, spot functional groups and draw isomers; at HL you also use spectra to identify a structure.
You need to be able to
- Draw and interpret full structural, condensed and skeletal formulas.
- Identify the common functional groups (alcohol, aldehyde, ketone, carboxylic acid, ester, amine, amide, halogenoalkane) and the homologous series they define.
- Name compounds with up to six carbons in the main chain using IUPAC rules, and classify alcohols and halogenoalkanes as primary, secondary or tertiary.
- Draw structural isomers and explain the boiling-point trend along a homologous series.
- HL: identify stereoisomers (cis–trans and optical), and deduce a structure from mass spectra, IR absorptions and ¹H NMR signals.
Key terms
- functional group
- homologous series
- structural isomer
- skeletal formula
- IUPAC nomenclature
- saturated
- unsaturated
- primary, secondary, tertiary
Common trap Not using the longest chain, or numbering from the wrong end. Find the longest continuous chain first (it may bend), then number from the end that gives the branch the lowest number: 2-methylpentane, never 4-methylpentane.
Practise S3.2 in Heptavo →Reactivity 1 — What drives chemical reactions?
The energy changes in reactions: measuring them, calculating them with cycles, burning fuels and, at HL, predicting spontaneity.
R1.1Measuring enthalpy changesSL
Exothermic and endothermic reactions, energy profiles and measuring enthalpy changes by calorimetry. Questions ask for q = mcΔT calculations, ΔH per mole and reasons why an experimental value differs from the literature value.
You need to be able to
- Identify exothermic and endothermic reactions from the temperature change, the sign of ΔH and energy profiles.
- Calculate the heat transferred with q = mcΔT using the mass of the solution, then divide by the moles of limiting reactant to get ΔH in kJ mol⁻¹.
- Correct for heat loss by extrapolating the cooling curve back to the time of mixing.
- Explain why measured enthalpies of combustion are less exothermic than data-book values, and suggest specific improvements.
- State the standard conditions used for enthalpy changes (100 kPa and 298 K).
Key terms
- system
- surroundings
- exothermic
- endothermic
- enthalpy change (ΔH)
- specific heat capacity
- calorimetry
- standard conditions
Common trap Getting the sign of ΔH wrong. q from mcΔT comes out positive; if the solution warmed, the reaction released heat and ΔH is negative, so you add the sign yourself.
Practise R1.1 in Heptavo →R1.2Energy cycles in reactionsSL
Hess's law and its use to find enthalpy changes that cannot be measured directly, from bond enthalpies and energy cycles. Questions are mostly calculations, plus explaining why bond-enthalpy answers are only approximate.
You need to be able to
- Calculate ΔH from average bond enthalpies: energy of bonds broken minus energy of bonds formed.
- Explain why bond-enthalpy answers are approximate: the values are averages over many compounds and apply to gases.
- Apply Hess's law to build and solve an energy cycle.
- HL: use standard enthalpies of formation or combustion in Hess's law calculations, and build a Born–Haber cycle to calculate a lattice enthalpy.
Key terms
- Hess's law
- state function
- average bond enthalpy
- energy cycle
- standard enthalpy of formation (HL)
- Born–Haber cycle (HL)
Common trap Mixing up the two subtraction orders. With bond enthalpies ΔH = bonds broken − bonds formed (reactants minus products); with formation data (HL) ΔH = ΣΔHf(products) − ΣΔHf(reactants). Draw the cycle if in doubt.
Practise R1.2 in Heptavo →R1.3Energy from fuelsSL
Combustion of organic fuels, the problems with fossil fuels, biofuels and fuel cells. Questions ask for balanced combustion equations, comparisons of fuels and evaluation of their environmental impact.
You need to be able to
- Write balanced equations for the complete and incomplete combustion of hydrocarbons and alcohols.
- Explain the hazards of incomplete combustion: toxic carbon monoxide and soot.
- Compare fossil fuels and biofuels, including what carbon neutral does and does not mean.
- Calculate specific energy (per gram) and energy density (per unit volume) from the enthalpy of combustion.
- Describe how a hydrogen fuel cell generates electricity and write the reaction at each electrode.
Key terms
- complete combustion
- incomplete combustion
- fossil fuel
- biofuel
- carbon neutral
- specific energy
- energy density
- fuel cell
Common trap Calling a biofuel carbon neutral without qualification. Burning it releases CO₂ just as a fossil fuel does; the claim rests on the crop having absorbed that CO₂ recently, and it ignores emissions from farming, processing and transport.
Practise R1.3 in Heptavo →R1.4Entropy and spontaneityHL
Entropy, Gibbs energy and how ΔH, ΔS and temperature together decide whether a reaction is spontaneous. Questions ask you to predict the sign of ΔS, calculate ΔG and find the temperature at which a reaction becomes spontaneous.
You need to be able to
- Predict the sign of ΔS from changes of state and changes in the number of moles of gas.
- Calculate ΔS⦵ for a reaction from standard entropy values.
- Calculate ΔG⦵ = ΔH⦵ − TΔS⦵, converting ΔS to kJ K⁻¹ mol⁻¹ first.
- Find the temperature at which ΔG = 0 and explain how temperature affects spontaneity for each combination of signs of ΔH and ΔS.
- Explain why a negative ΔG says nothing about how fast a reaction goes.
Key terms
- entropy (S)
- standard entropy change (ΔS⦵)
- Gibbs energy change (ΔG⦵)
- spontaneous reaction
- second law of thermodynamics
- kelvin
Common trap Leaving ΔS in J K⁻¹ mol⁻¹ while ΔH is in kJ mol⁻¹. Divide ΔS by 1000 before working out ΔH − TΔS, or the TΔS term is a thousand times too large.
Practise R1.4 in Heptavo →Reactivity 2 — How much, how fast and how far?
Reacting quantities, rates of reaction and chemical equilibrium.
R2.1How much? The amount of chemical changeSL
Stoichiometry: using balanced equations to calculate reacting amounts, limiting reactants, yields and titration results. Questions are multi-step calculations, often built on experimental data.
You need to be able to
- Balance equations with state symbols and use mole ratios to find reacting masses, gas volumes and solution volumes.
- Identify the limiting reactant by dividing each amount by its coefficient, and calculate the theoretical yield from it.
- Calculate percentage yield and atom economy, and explain what each one measures.
- Process titration data: select concordant titres, average them and calculate the unknown concentration.
- Use the molar volume of a gas at STP, 22.7 dm³ mol⁻¹ (273 K, 100 kPa).
Key terms
- stoichiometry
- limiting reactant
- excess reactant
- theoretical yield
- percentage yield
- atom economy
- titration
- concordant titres
Common trap Choosing the limiting reactant by comparing masses, or moles without the equation's ratio. Convert to moles, divide each by its coefficient, and the smallest result is limiting.
Practise R2.1 in Heptavo →R2.2How fast? The rate of chemical changeSL
Collision theory, the factors that affect rate, Maxwell–Boltzmann distributions and catalysis, with rate equations and mechanisms at HL. Questions ask you to explain rate changes, interpret rate graphs and sketch energy distributions.
You need to be able to
- Find a rate from the gradient of a concentration–time or volume–time graph, including a tangent at one point.
- Explain the effects of concentration, pressure, surface area, temperature and a catalyst using collision theory.
- Sketch Maxwell–Boltzmann curves at two temperatures with Ea marked, and use them to explain the effects of temperature and of a catalyst.
- Draw energy profiles with and without a catalyst, showing ΔH unchanged.
- HL: deduce orders and rate equations from data, and link the rate-determining step to a proposed mechanism.
Key terms
- rate of reaction
- collision theory
- activation energy (Ea)
- Maxwell–Boltzmann distribution
- catalyst
- transition state
- order of reaction (HL)
- rate constant (HL)
Common trap Saying a higher temperature speeds up a reaction mainly because there are more collisions. Collision frequency rises a little; the main reason is that a much larger fraction of collisions now have at least the activation energy.
Practise R2.2 in Heptavo →R2.3How far? The extent of chemical changeSL
Dynamic equilibrium, the equilibrium constant and Le Châtelier's principle, with K calculations and the reaction quotient at HL. Questions ask you to write Kc expressions and predict how changes affect the position of equilibrium.
You need to be able to
- Describe dynamic equilibrium: forward and reverse rates equal, concentrations constant but not necessarily equal.
- Write the Kc expression for a homogeneous equilibrium and interpret a large or small value.
- Predict how changes in concentration, pressure and temperature shift the position of equilibrium, and explain why a catalyst does not.
- Explain why only a change in temperature changes the value of Kc.
- HL: compare Q with K to predict the direction of change, calculate K from equilibrium data and relate ΔG⦵ to K.
Key terms
- dynamic equilibrium
- closed system
- equilibrium constant (Kc)
- position of equilibrium
- Le Châtelier's principle
- reaction quotient (Q)
- homogeneous equilibrium
Common trap Saying that adding a reactant increases Kc. The position shifts towards the products, but the concentrations readjust until the ratio is back at the same Kc; only temperature changes the constant.
Practise R2.3 in Heptavo →Reactivity 3 — What are the mechanisms of chemical change?
The four reaction types: proton transfer, electron transfer, radical reactions and reactions in which an electron pair is donated.
R3.1Proton transfer reactionsSL
Brønsted–Lowry acids and bases, pH, strong and weak acids and neutralisation, extended at HL to Ka, buffers and pH curves. Questions ask you to identify conjugate pairs, calculate pH and interpret titration data.
You need to be able to
- Identify Brønsted–Lowry acids and bases and their conjugate pairs in an equation, including amphiprotic species such as water.
- Calculate pH from [H⁺] and [H⁺] from pH, and use Kw = 1.00 × 10⁻¹⁴ at 298 K.
- Distinguish strong and weak acids, and describe tests that tell them apart at equal concentration (pH, conductivity, rate of reaction with magnesium).
- Write equations for neutralisation, and interpret the pH curve of a strong acid–strong base titration.
- HL: use Ka and pKa, interpret pH curves involving weak acids or bases, explain how buffers work and choose a suitable indicator.
Key terms
- Brønsted–Lowry acid
- Brønsted–Lowry base
- conjugate acid–base pair
- amphiprotic
- pH
- ionic product of water (Kw)
- strong acid
- weak acid
Common trap Confusing strength with concentration. A strong acid is fully dissociated at any concentration, so a dilute strong acid can have a higher pH than a concentrated weak one.
Practise R3.1 in Heptavo →R3.2Electron transfer reactionsSL
Oxidation states, redox equations, the activity series and electrochemical cells. Questions ask you to assign oxidation states, write half-equations, name oxidising and reducing agents and describe voltaic and electrolytic cells.
You need to be able to
- Assign oxidation states and use them to identify oxidation, reduction, the oxidising agent and the reducing agent.
- Balance redox half-equations in acidic solution using H₂O and H⁺, and combine them into an overall equation.
- Use an activity series to predict displacement reactions and which metals react with acids.
- Describe a voltaic cell (anode, cathode, salt bridge, direction of electron flow) and the electrolysis of a molten salt.
- Describe the oxidation of primary and secondary alcohols and the reduction of carbonyl compounds; HL adds standard electrode potentials and E⦵cell.
Key terms
- oxidation state
- oxidation
- reduction
- oxidising agent
- reducing agent
- half-equation
- anode
- cathode
Common trap Fixing the anode as positive. Oxidation always happens at the anode, but the anode is negative in a voltaic cell and positive in an electrolytic cell; learn the reaction at each electrode, not the sign.
Practise R3.2 in Heptavo →R3.3Electron sharing reactionsSL
Radicals and radical substitution, chiefly the reaction of alkanes with halogens in UV light. Questions ask you to write the initiation, propagation and termination steps and explain why a mixture of products forms.
You need to be able to
- Distinguish homolytic and heterolytic fission and show a radical's unpaired electron.
- Write the initiation step, two propagation steps and a termination step for the chlorination of methane.
- Explain why UV light is needed: it supplies enough energy to break the Cl–Cl bond homolytically.
- Explain why further substitution produces CH₂Cl₂, CHCl₃ and CCl₄ as well as CH₃Cl.
- Explain why alkanes are unreactive towards most reagents but react with radicals.
Key terms
- radical
- homolytic fission
- initiation
- propagation
- termination
- substitution reaction
- ultraviolet light
Common trap Writing a propagation step that makes a hydrogen radical, CH₄ + Cl• → CH₃Cl + H•. The chlorine radical removes a hydrogen atom: CH₄ + Cl• → CH₃• + HCl, then CH₃• + Cl₂ → CH₃Cl + Cl•.
Practise R3.3 in Heptavo →R3.4Electron-pair sharing reactionsSL
Reactions in which an electron pair is donated: nucleophilic substitution and electrophilic addition, extended at HL to Lewis acids and bases, SN1 and SN2 mechanisms and Markovnikov's rule. Questions ask you to classify reagents, predict products and, at HL, draw curly-arrow mechanisms.
You need to be able to
- Identify nucleophiles (electron-pair donors such as OH⁻, H₂O and NH₃) and electrophiles (electron-pair acceptors).
- Predict the products of nucleophilic substitution of halogenoalkanes by hydroxide ions.
- Predict the products of electrophilic addition of water, halogens and hydrogen halides to alkenes.
- Use bromine water to tell an alkene from an alkane, and explain the result.
- HL: draw SN1 and SN2 mechanisms, apply Markovnikov's rule and identify Lewis acids and bases in coordination bonds.
Key terms
- nucleophile
- electrophile
- heterolytic fission
- nucleophilic substitution
- electrophilic addition
- curly arrow
- Lewis acid
- Lewis base
Common trap Starting a curly arrow from an atom or a positive charge. Every curly arrow starts at an electron pair, either a lone pair or a bond, and points to where that pair ends up.
Practise R3.4 in Heptavo →Try 4 Heptavo questions
Real questions from the app, with the explanation after each answer.
S1.3 · Electron configurations
Which is the correct electron configuration of the S²⁻ ion? Sulfur has Z = 16.
- 1s² 2s² 2p⁶ 3s² 3p²
- 1s² 2s² 2p⁶ 3s² 3p⁴
- 1s² 2s² 2p⁶ 3s² 3p⁶
- 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
Show the answer and why
Answer: C
Sixteen electrons plus the two gained give eighteen, which fills the 3p sub-level and matches argon. The configuration ending 3p⁴ is the neutral sulfur atom, ending 3p² has lost two electrons rather than gained them, and adding a 4s pair takes the count to twenty.
S2.2 · The covalent model
Which sequence places the bonds in order of increasing length?
- C–C < C=C < C≡C
- C≡C < C=C < C–C
- They are all the same length
- C=C < C–C < C≡C
Show the answer and why
Answer: B
More shared pairs pull the nuclei closer, so a triple bond is the shortest and the strongest. Option A is the same list read as bond strength rather than length — the two run in opposite directions.
R1.1 · Measuring enthalpy changes
Which pair of quantities is not needed to calculate the heat change in a calorimetry experiment?
- the pressure of the atmosphere and the volume of the vessel
- the mass of solution and the temperature change
- the temperature change and the specific heat capacity
- the specific heat capacity and the mass of solution
Show the answer and why
Answer: A
The relationship q = mcΔT needs mass, specific heat capacity and temperature change, and nothing else. Atmospheric pressure matters to the definition of an enthalpy change, which is measured at constant pressure, but it never enters the arithmetic — that distinction between a condition and a variable is what this question tests.
R2.1 · How much? The amount of chemical change
What volume of 2.00 mol dm⁻³ hydrochloric acid is needed to prepare 250 cm³ of 0.150 mol dm⁻³ acid by dilution?
- 1.88 cm³
- 3.33 cm³
- 9.38 cm³
- 18.8 cm³
Show the answer and why
Answer: D
The dilute solution must contain 0.250 × 0.150 = 0.0375 mol, and that amount occupies 0.0375 ÷ 2.00 = 0.0188 dm³, or 18.8 cm³, of the concentrated acid. Dividing the concentrations the other way round gives 3.33 cm³, dividing by the concentration a second time gives 9.38 cm³, and 1.88 cm³ slips a factor of ten in the conversion from cubic decimetres.
