1Atomic Structure

1.1 Sub-atomic Particles

Atomic number Z = number of protons. Mass number A = protons + neutrons.

1.2 Electron Sub-shells and Orbitals

• Shells n = 1, 2, 3… subdivided into s, p, d sub-shells.
• s sub-shell: 1 orbital, max 2 electrons. p sub-shell: 3 orbitals, max 6 e⁻. d sub-shell: 5 orbitals, max 10 e⁻.
• Energy order: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p.
• Hund's rule: electrons occupy degenerate orbitals singly before pairing.
• Pauli exclusion: max 2 e⁻ per orbital, opposite spins.

1.3 Shapes of s and p Orbitals

The s orbital is spherical; the p orbital is dumb-bell shaped, with three orientations along x, y, z axes.

1.4 Ionisation Energy

Trends:

• ↑ across a period — nuclear charge increases, atomic radius decreases, shielding roughly constant.
• ↓ down a group — extra shell increases atomic radius and shielding, outweighs ↑ nuclear charge.
• Dip at Group 13 (e.g. Al < Mg) — the electron removed is from 3p, less penetrating than 3s.
• Dip at Group 16 (e.g. S < P) — spin-pair repulsion in doubly-occupied p orbital.

Successive IEs always increase. A large jump reveals movement to an inner shell — used to identify the group.

2Atoms, Molecules and Stoichiometry

2.1 Key Definitions

2.2 Mole Calculations

2.3 Yield and Limiting Reagent

Identify the limiting reagent by comparing mole ratios from the balanced equation; the one that gives the smallest product amount is limiting.

2.4 Ionic Equations

Spectator ions (those unchanged on both sides) are omitted. Atoms and charge must balance.

2.5 Common Ions to Memorise

NH₄⁺, OH⁻, NO₃⁻, CO₃²⁻, HCO₃⁻, SO₄²⁻, PO₄³⁻, MnO₄⁻, Cr₂O₇²⁻, Ag⁺, Zn²⁺, Pb²⁺.

3Chemical Bonding

3.1 Electronegativity

Increases across a period (↑ nuclear charge, ↓ radius); decreases down a group (↑ radius, more shielding). Large difference → ionic; small/zero → covalent. F is the most electronegative element.

3.2 Ionic Bonding

Examples: NaCl, MgO, CaF₂. Large difference in electronegativity required (typically > 1.7).

3.3 Covalent and Coordinate Bonding

• σ bond: head-on orbital overlap; first bond in a pair.
• π bond: sideways overlap of adjacent p orbitals; second/third bond in C=C, C≡C, C=O.
• Coordinate (dative) bond: both electrons supplied by one atom (e.g. NH₄⁺, Al₂Cl₆, NH₃·BF₃). Identical in strength once formed.

Period 3 and below can expand the octet (use d orbitals), giving SO₂, PCl₅, SF₆.

3.4 Metallic Bonding

Explains malleability, ductility, high mp/bp and electrical conductivity.

3.5 VSEPR Shapes

Lone pairs repel more than bonding pairs → reduce bond angle (NH₃ 107°, H₂O 104.5°).

3.6 Intermolecular Forces

• id-id (induced dipole - induced dipole / London): present in all molecules; strength ↑ with electron count / surface area.
• pd-pd (permanent dipole - permanent dipole): between polar molecules with net dipole moment.
• Hydrogen bonding: H bonded to N, O or F + lone pair on N/O/F of another molecule. Strongest intermolecular force.

4States of Matter

4.1 The Ideal Gas

Assumptions: negligible molecular volume, no intermolecular forces, elastic collisions, random motion. Real gases deviate at high pressure / low temperature.

4.2 Lattice Types

4.3 Allotropes of Carbon

• Diamond: sp³, four σ C–C bonds per atom, tetrahedral. Very hard, non-conductor.
• Graphite: sp², layers of hexagonal rings; delocalised π electrons → conducts along layers; layers slide → lubricant.
• Buckminsterfullerene (C₆₀): simple molecular sphere; held by id-id forces; low mp.

5Chemical Energetics (AS)

5.1 Enthalpy Change ΔH

Exothermic: ΔH < 0 (heat released; products lower than reactants). Endothermic: ΔH > 0. Standard conditions: 298 K, 101 kPa, ⦵.

5.2 Standard Enthalpy Definitions

5.3 Calorimetry

c for water = 4.18 J g⁻¹ K⁻¹. The negative sign converts "heat absorbed by water" to "enthalpy released by reaction".

5.4 Bond Energies

Breaking bonds is endothermic (+); making bonds is exothermic (−). Bond energies are averages, so calculated ΔH is approximate (exact only for gaseous diatomics).

5.5 Hess's Law

6Electrochemistry (AS)

6.1 Oxidation Number Rules

• Elements in standard state: 0.
• Combined O: −2 (except peroxides −1, F-oxides positive).
• Combined H: +1 (except metal hydrides −1).
• Sum of oxidation numbers = overall charge of species.

6.2 Oxidation and Reduction

Disproportionation: the same element is simultaneously oxidised and reduced (e.g. Cl₂ in NaOH).

6.3 Balancing Redox Equations

1. Write half-equations for oxidation and reduction.
2. Balance atoms (H with H⁺, O with H₂O, charge with e⁻).
3. Multiply to equalise electrons.
4. Add and cancel spectator species.

7Equilibria (AS)

7.1 Dynamic Equilibrium

Forward and reverse rates equal; concentrations of all species constant; requires a closed system. Approached from either side.

7.2 Le Chatelier's Principle

7.3 K_c and K_p

For aA + bB ⇌ cC + dD:

Partial pressure: p = mole fraction × total pressure. Only temperature changes K. Solids and pure liquids are omitted.

7.4 Brønsted–Lowry Acids and Bases

Conjugate pairs differ by one H⁺. Strong acids/bases fully dissociate; weak ones partially. Indicator choice: pKₐ of indicator must lie in the vertical region of the titration curve.

7.5 Industrial Processes

• Haber: N₂ + 3H₂ ⇌ 2NH₃ (ΔH negative). Compromise 400 °C, 200 atm, Fe catalyst.
• Contact: 2SO₂ + O₂ ⇌ 2SO₃, V₂O₅, 450 °C, 1–2 atm.

8Reaction Kinetics (AS)

8.1 Collision Theory

Molecules must collide with energy ≥ E_a and the correct orientation.

8.2 Boltzmann Distribution

Shows the distribution of molecular energies. Area to the right of E_a = fraction with enough energy.

• ↑ Temperature: curve flattens and shifts right → much larger fraction exceeds E_a → rate increases sharply.
• ↑ Concentration / pressure: more collisions per second.
• ↑ Surface area: more sites for collision.
• Catalyst: provides alternative pathway with lower E_a → larger fraction exceeds new E_a.

9Chemical Periodicity (Period 3)

9.1 Physical Trends Na → Ar

• Atomic radius decreases (↑ nuclear charge, ~constant shielding).
• Melting point: Na < Mg < Al (stronger metallic bond, more delocalised e⁻); Si highest (giant covalent); then drop to simple molecular P₄, S₈, Cl₂, Ar — S₈ > P₄ because of larger molecule.
• Electrical conductivity: Na < Mg < Al (more delocalised e⁻); Si semiconductor; non-metals do not conduct.

9.2 Elements with Oxygen and Water

Na, Mg, Al → ionic oxides; Si → giant covalent oxide; P, S → simple molecular acidic oxides.

9.3 Period 3 Chlorides with Water

NaCl (neutral, dissolves); MgCl₂ (slightly acidic); AlCl₃ (acidic hydrolysis); SiCl₄, PCl₅ (vigorous hydrolysis → HCl fumes + acidic oxoacid).

10Group 2

10.1 Reactivity Trends Mg → Ba

Reactivity increases down the group — atomic radius ↑ so 2 outer electrons more easily lost.

• + O₂ → MO (Ba forms peroxide BaO₂).
• + H₂O → M(OH)₂ + H₂ (vigour increases down group; Mg slow with cold water, fast with steam → MgO).
• + dilute acids → M²⁺ salt + H₂.

10.2 Solubility Trends

10.3 Thermal Stability

Stability of carbonates and nitrates increases down the group.

Explanation: larger cation has lower charge density → less polarising effect on the anion → harder to distort and decompose.

11Group 17 (Halogens)

11.1 Physical Properties Cl₂ → I₂

• Colour: Cl₂ pale green gas, Br₂ orange-brown liquid, I₂ grey-black solid (purple vapour).
• Volatility ↓ down group (stronger id-id forces as electron count increases).
• Bond energy Cl–Cl > Br–Br > I–I (longer bond → weaker overlap).

11.2 Halogens as Oxidising Agents

Strength: Cl₂ > Br₂ > I₂. Cl₂ displaces Br⁻ and I⁻; Br₂ displaces only I⁻.

11.3 Halide Ions as Reducing Agents

Strength: I⁻ > Br⁻ > Cl⁻. With conc. H₂SO₄:

• NaCl + H₂SO₄ → NaHSO₄ + HCl (steamy fumes; no further reduction).
• NaBr → HBr; HBr reduces some H₂SO₄ → Br₂ + SO₂.
• NaI → HI; HI reduces H₂SO₄ further → I₂, then H₂S, S (yellow solid, smell of rotten egg).

11.4 AgNO₃ Test for Halides

11.5 Disproportionation of Chlorine

Water purification: Cl₂ + H₂O ⇌ HCl + HOCl. HOCl kills bacteria.

12Nitrogen and Sulfur

12.1 Inertness of N₂

N≡N triple bond has bond energy 944 kJ mol⁻¹; non-polar. Requires high energy/catalyst to react.

12.2 Ammonia as a Base

NH₃ + H⁺ → NH₄⁺ (lone pair on N accepts proton; coordinate bond formed). NH₃ + H₂O ⇌ NH₄⁺ + OH⁻ (weak base).

12.3 NO_x in the Atmosphere

• Formed by lightning and in internal-combustion engines: N₂ + O₂ → 2NO at high temperature; 2NO + O₂ → 2NO₂.
• NO catalyses oxidation of SO₂: NO + ½O₂ → NO₂; NO₂ + SO₂ → NO + SO₃ → contributes to acid rain.
• Catalytic converter (Pt/Rh): 2NO + 2CO → N₂ + 2CO₂.

12.4 Acid Rain

SO₂ + H₂O → H₂SO₃; SO₃ + H₂O → H₂SO₄. Effects: damages plants, corrodes limestone and metal; treated by flue-gas desulfurisation with CaO/CaCO₃.

13Introduction to AS Organic Chemistry

13.1 Functional Groups (AS)

Alkene C=C, halogenoalkane C–X, alcohol –OH, aldehyde –CHO, ketone C=O, carboxylic acid –COOH, ester –COO–, amine –NH₂, nitrile –CN.

13.2 Nomenclature

IUPAC system. Longest chain → root (meth, eth, prop, but, pent, hex); add prefixes/suffixes; numbers indicate position; lowest locants overall.

13.3 Isomerism

• Structural: chain, positional, functional-group.
• Geometrical (cis/trans, E/Z): restricted rotation around C=C; two different groups on each sp² C.
• Optical: chiral C with four different groups → non-superimposable mirror images (enantiomers); rotate plane-polarised light in opposite directions.

13.4 Hybridisation

13.5 Reaction Types

Addition, substitution (free-radical, SN1, SN2), elimination, hydrolysis, condensation, oxidation, reduction.

13.6 Curly Arrows

Full arrow = movement of a pair of electrons (heterolytic). Half (fish-hook) arrow = movement of one electron (homolytic / radical). Arrow tail starts at a bond or lone pair; head points to where electrons go.

14Hydrocarbons

14.1 Alkanes

• Unreactive: non-polar, strong C–C and C–H bonds.
• Combustion: complete → CO₂ + H₂O; incomplete → CO + soot.
• Cracking: thermal (heat + Al₂O₃) → shorter alkanes + alkenes.

14.2 Free-Radical Substitution with X₂/UV

Three stages — write all equations using half-arrows:

• Initiation: Cl₂ → 2Cl• (UV)
• Propagation: CH₄ + Cl• → CH₃• + HCl; CH₃• + Cl₂ → CH₃Cl + Cl•
• Termination: Cl• + Cl• → Cl₂; CH₃• + Cl• → CH₃Cl; 2CH₃• → C₂H₆

14.3 Alkenes — Electrophilic Addition

14.4 Addition Polymerisation

n CH₂=CHX → −[CH₂–CHX]−_n. Examples: poly(ethene), PVC, polystyrene. Non-biodegradable; toxic on incineration (PVC → HCl).

15Halogen Compounds

15.1 Key Reactions

15.2 SN1 vs SN2

• SN1: two steps; rate = k[RX]; carbocation intermediate; tertiary preferred (more stable cation).
• SN2: one step; rate = k[RX][Nu⁻]; backside attack; primary preferred (less steric hindrance).
• Secondary: mixture.

15.3 C–X Reactivity

C–I (weakest, longest) > C–Br > C–Cl > C–F. AgNO₃/ethanol test: yellow AgI ppt forms fastest; AgCl white ppt slowest; AgF none.

16Hydroxy Compounds (Alcohols)

16.1 Classification

• Primary (1°): –CH₂OH (e.g. ethanol).
• Secondary (2°): –CHOH– (e.g. propan-2-ol).
• Tertiary (3°): –C(R)(R')OH (e.g. 2-methylpropan-2-ol).

16.2 Oxidation

Reagent: K₂Cr₂O₇(aq) / dil. H₂SO₄ (orange → green) or KMnO₄ / H⁺.

16.3 Other Reactions

• Na(s) → sodium alkoxide + ½H₂ (slow, less vigorous than water with Na).
• Conversion to halogenoalkane: HX(g); KCl + conc. H₂SO₄; PCl₃ + heat; PCl₅; SOCl₂.
• Dehydration: conc. H₂SO₄ at 170 °C, or Al₂O₃(s) hot → alkene.
• Esterification: + RCOOH / conc. H₂SO₄ cat. → ester + H₂O (reversible).

16.4 Iodoform Test

I₂ / NaOH(aq), warm. Yellow precipitate CHI₃ (triiodomethane) → presence of CH₃CH(OH)– or CH₃CO– group.

17Carbonyl Compounds

17.1 Common Reactions

17.2 Mechanism of HCN Addition (Nucleophilic Addition)

1. CN⁻ attacks δ⁺ carbon of C=O (curly arrow from CN⁻ lone pair to C).
2. π electrons of C=O move onto O → alkoxide intermediate.
3. H⁺ (from HCN or H₂O) protonates O → 2-hydroxynitrile.

18Carboxylic Acids and Esters

18.1 Carboxylic Acid Reactions

• + Reactive metals (Na, Mg) → salt + H₂.
• + Alkali → salt + H₂O.
• + Carbonate / hydrogencarbonate → salt + H₂O + CO₂ (test: effervescence).
• + Alcohol / conc. H₂SO₄ → ester + H₂O (reversible esterification).
• + LiAlH₄ → 1° alcohol (reduction).

18.2 Ester Hydrolysis

18.3 Acidity of Carboxylic Acids

Stronger than alcohols and phenol because the carboxylate ion is stabilised by delocalisation of negative charge over both O atoms (equivalent resonance structures).

19Nitrogen Compounds (AS)

19.1 Primary Amines

Production: R–X + excess NH₃ in ethanol, heated under pressure → R–NH₂. Multiple substitution can occur giving secondary/tertiary amines and quaternary salts.

19.2 Nitriles

• Production: R–X + KCN in ethanol, reflux → R–CN (chain lengthens by 1 C).
• Hydrolysis: H⁺/H₂O or OH⁻/H₂O, reflux → carboxylic acid (or its salt).
• Reduction: LiAlH₄ / dry ether → primary amine (R–CH₂NH₂).

19.3 Hydroxynitriles

Carbonyl + HCN / KCN(cat.) → 2-hydroxynitrile. Useful chain-extension and intermediate to α-hydroxy acids by hydrolysis.

20Polymerisation (AS)

20.1 Addition Polymerisation

n monomers join with no loss of atoms. Repeat unit: open up C=C and put brackets around the −[ ]−_n unit. Examples: poly(ethene), poly(chloroethene) (PVC), poly(propene), poly(phenylethene).

20.2 Environmental Issues

• Non-biodegradable — long C–C backbones.
• Combustion can release toxic gases: PVC → HCl; polynitriles → HCN.
• Solutions: recycling (mechanical, chemical), biodegradable alternatives, energy recovery.

21Organic Synthesis (AS)

21.1 Strategy

1. Identify start and target functional groups.
2. Plan a route via known intermediates (alcohol ↔ halogenoalkane ↔ alkene ↔ carbonyl ↔ acid ↔ ester).
3. Specify reagents, solvent, temperature and catalyst for each step.
4. Avoid steps that produce unwanted by-products; consider chain length.

21.2 Key Conversions

22Analytical Techniques (AS — IR & MS)

22.1 Infrared Spectroscopy

22.2 Mass Spectrometry

• M⁺ peak → molecular mass.
• [M+1]⁺ from ¹³C: n_C = 100 × I(M+1) / (1.1 × I(M)).
• [M+2]⁺: ~1:1 → Cl present (³⁵Cl/³⁷Cl); ~1:3 (actually 1:0.97) — never mind exact; ~ 1:1 → Br.
• Fragmentation: loss of 15 (CH₃), 17 (OH), 29 (CHO or C₂H₅), 31 (OCH₃), 43 (C₃H₇ or CH₃CO), 45 (COOH or OC₂H₅).

23Chemical Energetics (A2)

23.1 Lattice Energy

Magnitude (more negative) ↑ with higher charge and smaller radius (higher charge density) — e.g. MgO >> NaCl.

23.2 Born–Haber Cycle

Apply Hess's law:

• ΔH_at: enthalpy of atomisation (form 1 mol gaseous atoms).
• IE: ionisation energy (always +).
• EA: electron affinity (1st usually −; 2nd always +).

23.3 Enthalpy of Solution

ΔH_hyd: enthalpy when 1 mol of gaseous ions becomes hydrated. Always exothermic; more negative for smaller / higher-charge ions.

23.4 Entropy

ΔS positive for melting, boiling, dissolving, ↑ moles of gas, increased disorder. Calculate: ΔS = ΣS⦵(products) − ΣS⦵(reactants).

23.5 Gibbs Free Energy

Feasible (spontaneous) if ΔG < 0. Equilibrium temperature: T = ΔH⦵/ΔS⦵ (where ΔG = 0). Units: ΔH in kJ; convert ΔS to kJ K⁻¹.

24Electrochemistry (A2)

24.1 Electrolysis Calculations

24.2 Standard Electrode Potential E⦵

Measured relative to the Standard Hydrogen Electrode (Pt|H₂(101 kPa)|H⁺(1 mol dm⁻³)) at 298 K. Salt bridge: KNO₃(aq).

• More positive E⦵ → stronger oxidising agent (reduced more readily).
• More negative E⦵ → stronger reducing agent.

24.3 Cell EMF and Feasibility

Reaction feasible if E⦵_cell > 0. Electrons flow externally from more negative (anode, oxidised) to more positive (cathode, reduced).

24.4 Nernst Equation

↑ [oxidised] → E more positive; ↑ [reduced] → E more negative. Applies at 298 K.

25Equilibria (A2)

25.1 pH

25.2 Weak Acids

25.3 Buffers

Action: A⁻ + H⁺ → HA absorbs added H⁺; HA + OH⁻ → A⁻ + H₂O absorbs added OH⁻. Blood: H₂CO₃/HCO₃⁻ buffer keeps pH ~7.4.

25.4 Solubility Product

If ionic product > K_sp → precipitate forms. Common ion effect: adding a common ion reduces solubility.

25.5 Partition Coefficient K_pc

Constant at fixed T; depends on relative polarities. Used to model solvent extraction.

26Reaction Kinetics (A2)

26.1 Rate Equation

m, n found experimentally, not from stoichiometry. Overall order = m + n.

26.2 Orders

26.3 Rate-Determining Step

The slowest step in a mechanism. Only species in the rds (or before it) appear in the rate equation. Mechanism must be consistent with: overall equation; rate equation; rds slowness.

26.4 Catalysis

• Heterogeneous: different phase; adsorption → bond weakening → reaction → desorption. e.g. Fe in Haber, V₂O₅ in Contact, Pt/Rh in catalytic converters.
• Homogeneous: same phase; forms intermediate then regenerated. e.g. Fe²⁺/Fe³⁺ catalysing I⁻ + S₂O₈²⁻.

27Group 2 (A2 — Explanations)

27.1 Thermal Stability of Nitrates and Carbonates

Stability increases down the group. The cation's polarising power = charge / radius² ; smaller cations distort the anion's electron cloud, weakening the internal bond, and decomposition occurs at lower T.

27.2 Solubility Explained by ΔH_sol

ΔH_sol = −ΔH_latt + ΣΔH_hyd. For hydroxides, ΔH_latt falls faster than ΔH_hyd down the group → ΔH_sol more exothermic → more soluble. For sulfates, ΔH_hyd falls faster than ΔH_latt → less soluble.

28Chemistry of Transition Elements

28.1 Definition

28.2 Characteristic Properties

1. Variable oxidation states — 3d and 4s electrons are similar in energy.
2. Catalytic behaviour — vacant d orbitals; ability to change oxidation state.
3. Form complex ions — vacant d orbitals accept lone pairs from ligands.
4. Form coloured compounds — d-d electronic transitions absorb visible light.

28.3 Complex Ions and Ligands

• Ligand: a species with at least one lone pair donating to a metal ion via a dative bond.
• Monodentate: H₂O, NH₃, Cl⁻, CN⁻.
• Bidentate: en (1,2-diaminoethane), C₂O₄²⁻.
• Hexadentate: EDTA⁴⁻.
• Coordination number 6 → octahedral; 4 → tetrahedral or square planar; 2 → linear.

28.4 Colour — d-Orbital Splitting

In an octahedral field, the 5 d orbitals split into 2 higher (e_g) and 3 lower (t_2g) by energy ΔE. An electron absorbs a photon of energy hf = ΔE → complementary colour seen.

Ligand strength affects ΔE: I⁻ < Cl⁻ < F⁻ < H₂O < NH₃ < CN⁻.

28.5 Key Redox Reactions

28.6 Stability Constant K_stab

EDTA forms very stable 1:1 complexes (chelate effect): one EDTA replacing six H₂O ligands releases 6 mol of H₂O → ΔS positive → ΔG more negative.

28.7 Stereoisomerism in Complexes

• Cis/trans in square planar [Pt(NH₃)₂Cl₂]: cis is anti-cancer drug cisplatin; trans is inactive.
• Optical isomers in octahedral complexes with 3 bidentate ligands, e.g. [Ni(en)₃]²⁺.

29Introduction to A2 Organic Chemistry

29.1 New Functional Groups

Arene (benzene ring), halogenoarene, phenol, acyl chloride (RCOCl), amide (RCONH₂), amino acid (H₂N–CHR–COOH), azo (–N=N–).

29.2 Reactivity Patterns

• Benzene undergoes electrophilic substitution (aromatic stability preserved); alkenes undergo addition.
• Phenol is more reactive than benzene because the O lone pair donates into the ring.
• Acyl chlorides are far more reactive than carboxylic acids or esters towards nucleophiles.

30Arenes (Benzene Chemistry)

30.1 Structure of Benzene

Six sp² carbons; planar regular hexagon; delocalised π system above and below the ring (6 π electrons). All C–C bond lengths equal (~140 pm); intermediate between single (154) and double (134).

Evidence: enthalpy of hydrogenation of benzene is less exothermic than predicted for "1,3,5-cyclohexatriene" by ~150 kJ mol⁻¹ → aromatic stabilisation.

30.2 Electrophilic Substitution Mechanism

1. Form the electrophile.
2. Electrophile attacks the π system → curly arrow from ring to E⁺.
3. Arenium intermediate (carbocation with disrupted aromaticity, drawn as a partial circle with a + inside).
4. H⁺ leaves → aromaticity restored.

30.3 Key Reactions

30.4 Directing Effects

• 2,4-directors (activating): –OH, –NH₂, –NHR, –OR, –R (alkyl). Increase ring electron density.
• 3-directors (deactivating): –NO₂, –COOH, –CHO, –COR, –SO₃H. Withdraw electron density.

30.5 Side-chain Oxidation

Hot alkaline KMnO₄ then H⁺ → R–C₆H₅ → benzoic acid (regardless of side-chain length, provided it has at least one benzylic H).

31Halogenoarenes

31.1 Reactivity

Chlorobenzene does not readily react with aqueous NaOH (unlike chloropropane). Two reasons:

• Lone pair on Cl overlaps with ring π system → partial double-bond character to C–Cl, shortening and strengthening it.
• π electrons repel approaching nucleophile.

32Phenol

32.1 Acidity

Phenol is more acidic than ethanol because the phenoxide ion is stabilised by delocalisation of negative charge into the ring. Less acidic than carboxylic acid (where delocalisation is over two oxygens).

32.2 Reactions of the –OH Group

• + NaOH(aq) → sodium phenoxide + H₂O.
• + Na(s) → sodium phenoxide + H₂.
• Does not react with carbonates (less acidic than carbonic).

32.3 Ring Substitution

OH activates the ring strongly → phenol reacts with Br₂(aq) at room temperature without a catalyst → white ppt 2,4,6-tribromophenol.

32.4 Azo Coupling

Phenol + benzenediazonium chloride in NaOH(aq) below 10 °C → orange azo dye (–N=N– chromophore).

33Acyl Chlorides

33.1 Reactivity

Strongly polar C=O and good leaving group (Cl⁻) → most reactive carboxylic-acid derivative. Reacts vigorously with H₂O, alcohols, phenols and amines.

33.2 Key Reactions

33.3 Mechanism (Addition–Elimination)

1. Nucleophile attacks δ⁺ C of C=O.
2. π electrons of C=O move onto O (tetrahedral intermediate).
3. Lone pair on O reforms C=O, expelling Cl⁻.
4. Loss of H⁺ if necessary.

34Nitrogen Compounds (A2)

34.1 Basicity of Amines

Order: ethylamine > ammonia > phenylamine.

• Ethylamine: alkyl group is electron-donating (+I) → lone pair more available.
• Phenylamine: lone pair delocalised into ring → less available; weaker base.

34.2 Production of Amines

• RX + NH₃ in ethanol, heat, pressure → R–NH₂.
• LiAlH₄ on nitrile → R–CH₂NH₂; on amide → R–CH₂NH₂.
• Nitrobenzene + Sn / conc. HCl, reflux, then NaOH(aq) → phenylamine.

34.3 Diazonium Salts and Azo Dyes

Phenylamine + NaNO₂ / dil. HCl below 10 °C → C₆H₅N₂⁺Cl⁻ (benzenediazonium chloride). Reaction with phenol in NaOH(aq) → orange azo dye (R–N=N–R').

34.4 Amides

Weaker base than amines (N lone pair delocalised into C=O). Hydrolysis with dil. H⁺ or OH⁻ + heat → carboxylic acid + amine. LiAlH₄ reduction → amine.

34.5 Amino Acids

Have –NH₂ and –COOH on the same C. Form a zwitterion internally (–NH₃⁺ and –COO⁻).

• Low pH: –NH₃⁺ form → migrates to cathode.
• High pH: –COO⁻ form → migrates to anode.
• Peptide bond (–CO–NH–) formed by condensation between –COOH and –NH₂ of two amino acids.

35Condensation Polymerisation

35.1 Polyesters

Diol + dicarboxylic acid (or dioyl chloride) → polyester + H₂O (or HCl). Examples: Terylene (PET), made from benzene-1,4-dicarboxylic acid and ethane-1,2-diol.

35.2 Polyamides

Diamine + dicarboxylic acid → polyamide + H₂O. Examples: nylon-6,6 (hexane-1,6-diamine + hexanedioic acid); Kevlar (1,4-aromatic).

35.3 Proteins

Amino acid monomers join via peptide bonds → natural polyamides.

35.4 Biodegradability

Polyesters and polyamides contain –COO– and –CONH– bonds that can be hydrolysed by enzymes or acid/alkali → biodegradable. Addition polymers (C–C backbone) are not.

36Organic Synthesis (A2)

36.1 Strategy

1. Match functional groups via known transformations.
2. Consider regio-/stereoselectivity (e.g. directing effects in arenes; Markovnikov; chirality).
3. Order steps so reactive groups are installed/protected appropriately (e.g. nitration before reduction to amine; protect NH₂ before further nitration).
4. State every reagent + condition.

36.2 Chiral Drugs

Enantiomers can have very different biological activity (e.g. thalidomide). Industrial synthesis aims for single enantiomer using chiral catalysts or resolution from a racemic mixture.

37Analytical Techniques (A2 — TLC, GLC, NMR)

37.1 Thin-Layer Chromatography (TLC)

Stationary phase: thin layer of silica (or Al₂O₃) on a plate, polar. Mobile phase: non-polar solvent. Spots visualised under UV or with iodine.

More polar compound → greater interaction with stationary phase → lower R_f.

37.2 Gas–Liquid Chromatography (GLC)

Stationary phase: high-boiling-point non-polar liquid on a solid support. Mobile phase: unreactive carrier gas (N₂, He). Retention time identifies component; area under peak ∝ amount.

37.3 ¹³C NMR

Number of peaks = number of distinct C environments. Chemical shifts:

37.4 ¹H NMR

• Reference: TMS (δ = 0). Solvent: CDCl₃ (no H interfering).
• Each signal = different H environment.
• Integration ∝ number of H in that environment.
• n+1 rule: n equivalent H on adjacent C → n+1 peaks (singlet, doublet, triplet, quartet…).
• D₂O shake: O–H and N–H protons exchange → peak disappears.

📋Data, Constants and Exam Information

Fundamental Constants

Paper Structure

Command Words

Paper 5 — Linearisation Cheatsheet

Cation Tests (Paper 3 — Qualitative Analysis)

Anion & Gas Tests

Exam Tips