NJC J2 H2 Chemistry, 2024 A-Level – Topics Not Yet Tested in Paper 2
Difficulty: Mixed (Foundation to Advanced) | Prerequisites: Varies by topic; assumes basic understanding of atomic structure, bonding, and organic functional groups.
Big picture: After Paper 2, the NJC analysis identifies a wide range of topics that have not yet appeared. These span Physical Chemistry (atomic structure, bonding, energetics calculations, kinetics, electrochemistry, equilibrium, ionic equilibrium), Inorganic Chemistry (Period 3, Group 2, Group 17, transition metals), and Organic Chemistry (mechanisms, isomerism, directing effects, distinguishing tests). These topics remain fair game for Papers 1 and 3. The NJC analysis specifically warns that tested topics can reappear in other papers, so treat this as a "high priority" revision list, not a guarantee of what will appear.
This document covers every topic flagged as untested after Paper 2. Physical Chemistry topics include ionisation energy trends, VSEPR, Hess's law variants, rate equations, electrochemistry, and equilibrium calculations. Inorganic Chemistry covers Period 3, Group 2 and 17 properties, and transition metal chemistry. Organic Chemistry focuses on mechanisms, directing effects, hybridisation, and distinguishing tests. Use these notes as a targeted revision checklist for Papers 1 and 3.
Ionisation energy (IE)
The energy required to remove one mole of electrons from one mole of gaseous atoms (or ions) to form gaseous cations. In simple terms, how tightly an atom holds onto its electrons.
VSEPR theory
Valence Shell Electron Pair Repulsion theory. Electron pairs around a central atom repel each other and arrange themselves to minimise repulsion, determining molecular shape.
Bond energy (BE)
The average energy required to break one mole of a particular covalent bond in the gaseous state. Used in Hess's law: ΔH_r = Σ BE(bonds broken) – Σ BE(bonds formed).
Gibbs Free Energy (ΔG)
ΔG = ΔH – TΔS. A reaction is spontaneous when ΔG < 0. Combines enthalpy and entropy to predict feasibility.
Rate equation
An expression relating the rate of reaction to the concentrations of reactants raised to their orders: rate = k[A]^m[B]^n. The orders (m, n) are found experimentally, not from the stoichiometric equation.
Standard electrode potential (E⁰)
The potential of a half-cell measured against the standard hydrogen electrode under standard conditions. Used to predict whether a redox reaction is feasible: E⁰_cell = E⁰(cathode) – E⁰(anode).
Kc and Kp
Equilibrium constants in terms of concentration (Kc) and partial pressure (Kp). They are constant at a given temperature and shift only when temperature changes.
Lewis acid / Lewis base
A Lewis acid is an electron-pair acceptor; a Lewis base is an electron-pair donor. This is the broadest definition of acid-base behaviour.
Electrolysis
Using an external electrical current to drive a non-spontaneous redox reaction. Cations are reduced at the cathode; anions are oxidised at the anode.
Directing effects (EDG / EWG)
In aromatic substitution, electron-donating groups (EDG) direct incoming electrophiles to the 2,4 positions (ortho/para). Electron-withdrawing groups (EWG) direct to the 3 position (meta).
Hybridisation
The mixing of atomic orbitals to form new hybrid orbitals (sp, sp², sp³) with equivalent energy and shape. Determines bond angles and molecular geometry.
Complex ion
A central metal ion surrounded by ligands (molecules or ions that donate lone pairs). Defined by the metal, its oxidation state, the ligands, and the coordination number.
Ionisation energy trends across Period 3: general increase left to right (increasing nuclear charge, same shielding). Two exceptions:
Group 2 → Group 13 drop: the 3p electron in Al is easier to remove than the 3s² pair in Mg (less penetration, more shielding by 3s²).
Group 15 → Group 16 drop: the paired 3p electron in S experiences inter-electron repulsion, making it easier to remove than an unpaired 3p electron in P.
Down a group: IE decreases (more shells, greater shielding, weaker nuclear attraction on outermost electron).
Successive IE: large jumps indicate removal of an electron from a new, inner shell.
Shapes of orbitals: s is spherical, p is dumbbell-shaped (three orientations), d has more complex shapes (cloverleaf for most).
Deflection of charged particles in an electric field: protons deflect toward the negative plate, electrons deflect toward the positive plate (more sharply, due to lower mass), neutrons pass straight through.
Dot-and-cross diagrams: show valence electrons, with dots for one atom and crosses for another. Include lone pairs.
VSEPR: count electron pairs (bonding + lone) around the central atom. Lone pairs repel more than bonding pairs. Shapes: linear (2), trigonal planar (3), tetrahedral (4), trigonal bipyramidal (5), octahedral (6).
Structures to know: diamond (giant covalent, tetrahedral, very hard), graphite (giant covalent layers, delocalised electrons, conducts electricity, slippery), ice (hydrogen-bonded open structure, less dense than liquid water), copper (metallic bonding, close-packed), iodine (simple molecular, held by weak London forces, low melting point).
Orbital overlap diagrams: sigma bonds from head-on overlap, pi bonds from side-on overlap of p orbitals.
Melting/boiling points: determined by the type and strength of bonding/intermolecular forces to overcome.
ΔH_r = (±) mcΔT / n: calorimetry calculation. m = mass of solution, c = specific heat capacity (usually 4.18 J g⁻¹ K⁻¹), ΔT = temperature change, n = limiting moles.
Hess's law routes:
ΔH_r = Σ ΔH_c(reactants) – Σ ΔH_c(products)
ΔH_r = Σ ΔH_f(products) – Σ ΔH_f(reactants)
ΔH_r = Σ BE(bonds broken) – Σ BE(bonds formed)
ΔG = ΔH – TΔS: predicts spontaneity. Exothermic + entropy increase = always spontaneous. Endothermic + entropy decrease = never spontaneous. Mixed signs: depends on temperature.
Thermochemical equations: include state symbols and ΔH value. Coefficients refer to moles.
Order of reaction: the power to which a reactant's concentration is raised in the rate equation. Found experimentally (initial rates method, half-life method).
Half-life: constant half-life = first order. For first order: t₁/₂ = ln2 / k.
Rate equation: rate = k[A]^m[B]^n. Overall order = m + n. Units of k depend on overall order.
Catalysis: homogeneous (same phase as reactants), heterogeneous (different phase, often solid catalyst with gaseous/liquid reactants). Enzymes: biological catalysts, substrate binds to active site, saturation at high [substrate].
Maxwell-Boltzmann distribution: shows distribution of molecular kinetic energies. Higher temperature shifts the curve right and flattens it; a catalyst lowers the activation energy but does not change the distribution.
Energy profile diagrams: show E_a (activation energy), ΔH (enthalpy change), and the effect of a catalyst (lower E_a, same ΔH).
Oxidation numbers: assign using rules (free element = 0, monatomic ion = charge, O = –2 usually, H = +1 usually).
Balancing redox equations: split into half-equations, balance atoms and charges, combine.
Electrochemical cell: two half-cells connected by a salt bridge. E⁰_cell = E⁰(cathode) – E⁰(anode). Positive E⁰_cell means the reaction is feasible under standard conditions.
ΔG = –nFE⁰_cell (n = moles of electrons transferred, F = Faraday constant = 96500 C mol⁻¹).
Electrolysis: at the cathode, cations are reduced (preferentially the one with the more positive E⁰). At the anode, anions are oxidised (preferentially the one with the more negative E⁰).
Applications: purification of copper (impure Cu anode, pure Cu cathode, CuSO₄ electrolyte), anodising of aluminium.
Kc = products over reactants, each raised to their stoichiometric coefficients. Only includes aqueous and gaseous species.
Kp = partial pressures of products over reactants. Partial pressure = mole fraction × total pressure.
Le Chatelier's principle: a system at equilibrium responds to oppose a change (concentration, pressure, temperature). Only temperature changes K.
Arrhenius: acid produces H⁺ in water, base produces OH⁻.
Bronsted-Lowry: acid is a proton donor, base is a proton acceptor.
Lewis: acid is an electron-pair acceptor, base is an electron-pair donor.
pH of strong acids/bases: straightforward [H⁺] or [OH⁻] calculation.
pH of weak acids: use Ka and ICE table. pH = –lg[H⁺].
pH of salts: consider hydrolysis of the cation or anion.
Buffers: weak acid + conjugate base (or weak base + conjugate acid). pH ≈ pKa + lg([A⁻]/[HA]).
Titration curves: know the shapes for strong/strong, strong/weak, weak/strong, weak/weak. Identify equivalence point, buffer region, half-equivalence point (pH = pKa).
Physical properties: atomic/ionic radius decreases across the period (increasing nuclear charge). Melting points: Na, Mg, Al increase (metallic bonding strengthens with more delocalised electrons and smaller ionic radius); Si is very high (giant covalent); P₄, S₈, Cl₂, Ar are low (simple molecular, weak intermolecular forces).
Reactions of chlorides with water: NaCl and MgCl₂ dissolve (ionic). AlCl₃ hydrolyses partially (produces acidic solution). SiCl₄ hydrolyses vigorously in water to give SiO₂ and HCl (white fumes). PCl₃ and PCl₅ hydrolyse to give phosphoric acid and HCl.
Oxides: Na₂O and MgO are basic (ionic). Al₂O₃ is amphoteric. SiO₂ is weakly acidic (giant covalent). P₄O₁₀, SO₂, SO₃, Cl₂O₇ are acidic (covalent).
Reactivity increases down the group: Be < Mg < Ca < Sr < Ba. Ionisation energies decrease, so it becomes easier to lose the two outer electrons.
Reactions with water: Mg reacts slowly with cold water (forms Mg(OH)₂), reacts vigorously with steam. Ca, Sr, Ba react increasingly vigorously with cold water, forming M(OH)₂ and H₂.
Oxidising strength decreases down the group: F₂ > Cl₂ > Br₂ > I₂. A higher halogen can displace a lower one from solution.
Volatility increases up the group (weaker London forces for smaller molecules).
Thermal decomposition of hydrogen halides: HF is very stable (strong H–F bond). HI decomposes most easily (weakest H–I bond). Stability: HF > HCl > HBr > HI.
Reaction with Ag⁺(aq): AgF is soluble (no precipitate). AgCl is white, soluble in dilute NH₃. AgBr is cream, soluble in concentrated NH₃. AgI is yellow, insoluble in NH₃.
Physical properties: high density, high melting points, hard, good conductors. Due to strong metallic bonding from 3d and 4s electrons.
Complex ions: a central metal ion with ligands attached by dative (coordinate) bonds. Coordination number = total number of dative bonds to the metal.
Ligands: monodentate (one lone pair donated, e.g. H₂O, NH₃, Cl⁻), bidentate (two, e.g. en, C₂O₄²⁻), hexadentate (six, e.g. EDTA⁴⁻).
Catalysis: homogeneous (same phase, e.g. Fe²⁺/Fe³⁺ in the reaction of S₂O₈²⁻ with I⁻) and heterogeneous (different phase, e.g. Fe in Haber process, V₂O₅ in Contact process).
Chemistry of Cu: Cu²⁺(aq) is blue, [Cu(H₂O)₆]²⁺. Addition of NaOH gives blue Cu(OH)₂ precipitate. Addition of excess NH₃ gives deep blue [Cu(NH₃)₄(H₂O)₂]²⁺.
Chemistry of Cr: Cr³⁺(aq) is green/violet, [Cr(H₂O)₆]³⁺. Addition of NaOH gives grey-green Cr(OH)₃. Excess NaOH gives green [Cr(OH)₆]³⁻ (amphoteric). CrO₄²⁻ (yellow, alkaline) ⇌ Cr₂O₇²⁻ (orange, acidic).
Chemistry of Fe: Fe²⁺ gives green solution, pale green Fe(OH)₂ with NaOH. Fe³⁺ gives yellow/brown solution, brown Fe(OH)₃ with NaOH.
sp³: four sigma bonds / lone pairs, tetrahedral, 109.5°. Example: CH₄, saturated carbons.
sp²: three sigma bonds + one pi bond, trigonal planar, 120°. Example: C=C in alkenes, C in benzene.
sp: two sigma bonds + two pi bonds (or one sigma + one pi + lone pairs), linear, 180°. Example: C≡C in alkynes, CO₂.
Draw labelled diagrams showing the hybrid orbitals and the unhybridised p orbitals forming pi bonds.
Acidity of carboxylic acids vs. alcohols vs. phenols: carboxylic acids are the most acidic (carboxylate anion is stabilised by delocalisation over two equivalent oxygens). Phenols are weakly acidic (phenoxide stabilised by delocalisation into the ring). Alcohols are the least acidic.
Basicity of amines: aliphatic amines are more basic than NH₃ (electron-donating alkyl groups increase electron density on N). Aromatic amines (e.g. phenylamine) are less basic than NH₃ (lone pair on N is delocalised into the ring, less available for protonation).
Structural isomerism: same molecular formula, different structural formula (chain, position, functional group isomers).
Geometric (cis-trans / E-Z) isomerism: restricted rotation around a C=C double bond with two different groups on each carbon.
Electron-donating groups (EDG): –OH, –NH₂, –OR, –alkyl. Activate the ring (make substitution faster) and direct to ortho/para (2,4) positions.
Electron-withdrawing groups (EWG): –NO₂, –COOH, –CHO, –COR. Deactivate the ring and direct to meta (3) position.
Mechanism: electrophilic aromatic substitution. Draw the carbocation intermediate and show how EDG stabilises it at ortho/para positions through resonance.
Delocalisation of electrons (pi or lone pair) over multiple atoms stabilises a structure by spreading charge.
Examples: benzene (6 p electrons delocalised), carboxylate ion (charge spread over two oxygens), phenoxide ion, amide bond.
In the addition of HX to an unsymmetrical alkene, the H adds to the carbon with more hydrogen atoms already attached (the less substituted carbon), and the X adds to the more substituted carbon.
Reason: the more substituted carbocation intermediate is more stable (hyperconjugation and inductive effects).
These are problem-solving questions where you identify unknown compounds from their reactions.
Key tests to memorise:
Br₂(aq): decolourises with alkenes (addition) and phenols (substitution).
2,4-DNPH: forms orange precipitate with aldehydes and ketones.
Tollens' reagent (silver mirror test): positive with aldehydes only.
Fehling's solution: positive with aldehydes only (brick-red precipitate).
FeCl₃: purple/violet colouration with phenols.
Na₂CO₃: effervescence (CO₂) with carboxylic acids.
Na metal: effervescence (H₂) with alcohols, carboxylic acids, phenols (any O–H or N–H).
Free radical substitution: initiation (homolytic fission by UV), propagation (chain), termination.
Electrophilic addition: to alkenes (e.g. HBr, Br₂). Draw the carbocation intermediate.
Electrophilic aromatic substitution: nitration (HNO₃/H₂SO₄), halogenation (Cl₂/AlCl₃), Friedel-Crafts acylation.
Nucleophilic addition: to carbonyl group (e.g. HCN/KCN to aldehydes/ketones).
SN1: two-step, carbocation intermediate, racemisation. Favoured by tertiary substrates.
Elimination: E1 and E2 to form alkenes from halogenoalkanes using hot ethanolic NaOH.
Students often think that IE always increases smoothly across a period. It does not. The two dips (Group 2→13 and Group 15→16) are tested regularly and must be explained in terms of subshell and electron pairing.
A frequent mistake in VSEPR is treating all electron pairs as equivalent. Lone pairs occupy more space than bonding pairs, so bond angles are compressed (e.g. H₂O is 104.5°, not 109.5°).
Students sometimes believe that a catalyst shifts the position of equilibrium. It does not. A catalyst increases the rate of both forward and reverse reactions equally, so equilibrium is reached faster but Kc and Kp are unchanged.
In electrolysis, students sometimes reverse the electrode labels. Remember: cations go to the cathode (reduction), anions go to the anode (oxidation). "CATions at the CAThode."
A common organic error is confusing EDG and EWG directing effects. EDG directs ortho/para, EWG directs meta. Students sometimes assume all substituents activate the ring; nitro groups and carbonyls deactivate it.
⚠️ The NJC analysis warns that topics tested in Paper 2 can still reappear in Papers 1 and 3. Do not skip any topic.
⚠️ Organic mechanisms (free radical substitution, electrophilic addition, electrophilic aromatic substitution, nucleophilic addition) are heavily tested in Paper 3. Draw every curly arrow, lone pair, and intermediate clearly.
⚠️ Hybridisation with labelled diagrams is a common extended-response question. Practise drawing sp, sp², sp³ with the unhybridised p orbitals.
⚠️ Kinetics calculations (rate equation from experimental data, graph plotting for order determination) are classic Paper 1 MCQ material.
⚠️ Transition metal colour, ligand exchange, and catalysis are high-frequency topics across all papers.
⚠️ Ionic equilibrium (buffer calculations, titration curves, pH of salts) is one of the most mark-dense areas. Know the Henderson-Hasselbalch equation and how to sketch all four titration curve shapes.
⚠️ Distinguishing tests in organic chemistry: memorise the reagent, conditions, observation, and conclusion for each functional group test. These appear in both structured and free-response questions.
True or false: Across Period 3, ionisation energy always increases from Na to Ar with no exceptions. (False – there are dips from Mg to Al and from P to S.)
Fill in the blank: SiCl₄ reacts vigorously with water to form ______ and HCl. (SiO₂)
True or false: A catalyst changes the value of the equilibrium constant. (False – a catalyst does not change Kc or Kp.)
Fill in the blank: In electrophilic aromatic substitution, an EDG directs to the ______ and ______ positions. (ortho and para / 2 and 4)
True or false: AgBr is soluble in dilute ammonia. (False – AgBr requires concentrated NH₃ to dissolve. AgCl dissolves in dilute NH₃.)
Q: Explain why the first ionisation energy of aluminium is lower than that of magnesium.
A: Al's outer electron is in the 3p subshell, which is higher in energy and further from the nucleus than the 3s subshell of Mg. The 3p electron is also shielded by the 3s² electrons. Therefore, less energy is required to remove it.
Q: Predict the shape and bond angle of PCl₄⁺.
A: P has 4 bonding pairs and 0 lone pairs around it in PCl₄⁺. Shape: tetrahedral. Bond angle: 109.5°.
Q: The rate equation for a reaction is rate = k[A]²[B]. If [A] is doubled and [B] is halved, what happens to the rate?
A: New rate = k(2[A])²(0.5[B]) = k × 4[A]² × 0.5[B] = 2k[A]²[B]. The rate doubles.
Q: Explain why phenylamine is a weaker base than ethylamine.
A: In phenylamine, the lone pair on nitrogen is delocalised into the benzene ring. This makes it less available for accepting a proton. In ethylamine, the ethyl group is electron-donating (by induction), which increases electron density on the nitrogen lone pair, making it a better proton acceptor.
Q: Draw the mechanism for the nitration of benzene.
A: Step 1 – Generate the electrophile: HNO₃ + H₂SO₄ → NO₂⁺ + HSO₄⁻ + H₂O. Step 2 – Electrophilic attack: the delocalised pi electrons of benzene attack NO₂⁺, forming a positively charged intermediate (with a broken ring). Step 3 – Loss of H⁺: a base (HSO₄⁻) removes H⁺ from the intermediate, restoring aromaticity. Product: nitrobenzene.
Atomic Structure underpins everything: IE trends explain Group 2 reactivity and Period 3 properties. Chemical Bonding connects to Organic Chemistry (hybridisation, sigma/pi bonds) and to the structures of Period 3 chlorides and oxides. Energetics (Hess's law, Gibbs Free Energy) links to Electrochemistry (ΔG = –nFE⁰) and to Ionic Equilibrium (buffer calculations require Ka, which is a thermodynamic quantity). Kinetics connects to Transition Metals (catalysis) and to Organic Chemistry (mechanism determines rate law). Organic mechanisms tie together functional group chemistry, directing effects, and distinguishing tests into the deductive problems that dominate Paper 3.
Ionisation energy, IE trends, Period 3, subshell, VSEPR, molecular shape, dot-and-cross, diamond, graphite, ice, metallic bonding, Hess's law, calorimetry, bond energy, Gibbs Free Energy, spontaneity, rate equation, order of reaction, half-life, Maxwell-Boltzmann, catalyst, homogeneous, heterogeneous, enzyme, oxidation number, redox, electrochemical cell, E⁰, electrolysis, Faraday, Kc, Kp, Le Chatelier, Arrhenius, Bronsted-Lowry, Lewis acid, Lewis base, pH, buffer, Henderson-Hasselbalch, titration curve, Period 3 chlorides, Period 3 oxides, Group 2, Group 17, halogens, AgNO₃ test, transition metal, complex ion, ligand, coordination number, Cu, Cr, Fe, hybridisation, sp sp² sp³, EDG, EWG, directing effect, Markovnikov, resonance, distinguishing test, organic mechanism, H2 Chemistry, A-Level, NJC, 2024