9701 Chemistry — Notes & AI Tutor

These notes are AI-assisted study material. Always cross-check against the official 9701 syllabus or your teacher before relying on them in an exam.

1Atomic Structure

Sub-atomic particles, isotopes, electron configuration in shells/sub-shells/orbitals, and trends in ionisation energy.

1.1 Sub-atomic Particles

Particle	Relative charge	Relative mass
Proton	+1	1
Neutron	0	1
Electron	−1	1/1836

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

Isotopes: atoms of the same element with the same number of protons but different numbers of neutrons. Same chemical properties (same electron configuration); different physical properties (different mass and density).

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.

Cr is [Ar] 3d⁵ 4s¹ and Cu is [Ar] 3d¹⁰ 4s¹ — exceptions for extra stability of half-filled / full d sub-shell.

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

First IE: energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions. X(g) → X⁺(g) + e⁻

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.

Worked example

Successive IEs of an element (kJ mol⁻¹): 738, 1451, 7733, 10 540. The big jump between IE₂ and IE₃ means there are only 2 electrons in the outer shell → element is in Group 2.

2Atoms, Molecules and Stoichiometry

The mole concept, relative masses, empirical and molecular formulae, and calculation of reacting quantities.

2.1 Key Definitions

Mole: the amount of substance that contains as many particles as there are atoms in 12 g of carbon-12. L = 6.022 × 10²³ mol⁻¹.

Relative atomic mass A_r: weighted mean mass of an atom of an element, relative to 1/12 of the mass of a ¹²C atom.

Empirical formula: simplest whole-number ratio of atoms. Molecular formula: the actual number of atoms in a molecule.

2.2 Mole Calculations

moles n = mass / M_r

Solutionsmoles = concentration × volume / 1000 (if V is in cm³)

Gasesn = V / 24.0 dm³ (room conditions); V / 22.4 dm³ at s.t.p.

pV = nRT (R = 8.31 J K⁻¹ mol⁻¹)

2.3 Yield and Limiting Reagent

% yield = (actual moles / theoretical moles) × 100%

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.

Example

Full: AgNO₃ + NaCl → AgCl + NaNO₃
Ionic: Ag⁺ + Cl⁻ → AgCl(s)

2.5 Common Ions to Memorise

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

3Chemical Bonding

Electronegativity, ionic/covalent/metallic bonding, VSEPR shapes, and intermolecular forces (including hydrogen bonding).

3.1 Electronegativity

Electronegativity: the power of an atom in a covalent bond to attract the shared pair of electrons to itself.

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

Ionic bond: the electrostatic attraction between oppositely charged ions formed by electron transfer.

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

3.3 Covalent and Coordinate Bonding

Covalent bond: the electrostatic attraction between the nuclei of two atoms and the shared pair of electrons between them.

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

Metallic bond: the electrostatic attraction between the lattice of positive metal ions and the delocalised sea of electrons.

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

3.5 VSEPR Shapes

Molecule	Shape	Bond angle
BF₃	Trigonal planar	120°
CO₂	Linear	180°
CH₄	Tetrahedral	109.5°
NH₃	Pyramidal	107°
H₂O	Bent / non-linear	104.5°
PF₅	Trigonal bipyramidal	90° & 120°
SF₆	Octahedral	90°

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.

H bonding explains why ice is less dense than water (open tetrahedral lattice), and the high bp of H₂O, NH₃ and HF compared to other Group 16/15/17 hydrides.

Common mistake: calling H-bond "an actual bond between molecules". Be precise: it is an intermolecular force, weaker than a covalent bond but stronger than other intermolecular forces.

4States of Matter

Ideal gas behaviour and the four lattice types (ionic, simple molecular, giant covalent, metallic).

4.1 The Ideal Gas

pV = nRT

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

4.2 Lattice Types

Type	Example	mp/bp	Conductivity
Giant ionic	NaCl, MgO	High	Only molten/aqueous
Simple molecular	I₂, C₆₀, ice	Low	None
Giant covalent	SiO₂, diamond, graphite	Very high	None (graphite conducts)
Giant metallic	Cu, Mg	High	Solid & molten

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)

Enthalpy changes, standard definitions, calorimetry, Hess's law and bond energies.

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

ΔH_f⦵ (formation): enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions.
ΔH_c⦵ (combustion): enthalpy change when 1 mole of a substance is completely burned in oxygen under standard conditions.
ΔH_neut⦵: enthalpy change when 1 mole of water is formed by the neutralisation of an acid by an alkali under standard conditions.
ΔH_r⦵ (reaction): enthalpy change for a reaction as the equation is written under standard conditions.

5.3 Calorimetry

q = mcΔT ⟹ ΔH = −mcΔT / n

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

ΔH = Σ(bonds broken) − Σ(bonds formed)

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

Hess's Law: the enthalpy change of a reaction is independent of the route, depending only on the initial and final states.

Using ΔH_f: ΔH_r = ΣΔH_f(products) − ΣΔH_f(reactants). Using ΔH_c: ΔH_r = ΣΔH_c(reactants) − ΣΔH_c(products) — opposite direction.

6Electrochemistry (AS)

Oxidation numbers, half-equations and balancing redox processes.

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

Oxidation: loss of electrons / increase in oxidation number. Reduction: gain of electrons / decrease in oxidation number. (OIL RIG)

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

6.3 Balancing Redox Equations

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

Worked example

MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O

7Equilibria (AS)

Dynamic equilibrium, Le Chatelier's principle, K_c and K_p, Brønsted–Lowry acids and bases, and industrial processes.

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

When a system at equilibrium is disturbed, it shifts in the direction that minimises the disturbance.

Disturbance	Position shifts to…
↑ reactant concentration	products (right)
↑ temperature	endothermic side
↑ pressure	side with fewer moles of gas
Catalyst	no shift; equilibrium reached faster

7.3 K_c and K_p

For aA + bB ⇌ cC + dD:

K_c = [C]^c[D]^d / ([A]^a[B]^b)

K_p = (p_C^c·p_D^d) / (p_A^a·p_B^b) (partial pressures)

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

7.4 Brønsted–Lowry Acids and Bases

Acid: a proton (H⁺) donor. Base: a proton acceptor.

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.

Always state that conditions are a compromise between yield and rate.

8Reaction Kinetics (AS)

Collision theory, activation energy, Boltzmann distributions, and effect of catalysts.

8.1 Collision Theory

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

Activation energy E_a: the minimum energy reacting particles must possess for a successful collision.

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.

Common error: describing temperature effect as "molecules move faster so they collide more often" — this is true but minor. The dominant factor is the exponential increase in the fraction with E ≥ E_a.

9Chemical Periodicity (Period 3)

Trends in Period 3: atomic radius, melting point, conductivity, and reactions of elements, oxides and chlorides.

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.

Element	Oxide formula	Reaction with water	pH
Na	Na₂O	Na₂O + H₂O → 2NaOH	~13
Mg	MgO	Slow; Mg(OH)₂ slightly soluble	~9
Al	Al₂O₃	Insoluble; amphoteric	~7
Si	SiO₂	No reaction	—
P	P₄O₁₀	→ H₃PO₄	~2
S	SO₂ / SO₃	→ H₂SO₃ / H₂SO₄	<2

9.3 Period 3 Chlorides with Water

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

SiCl₄ + 2H₂O → SiO₂ + 4HCl

10Group 2

Trends in reactivity, solubility of hydroxides and sulfates, and thermal stability of nitrates/carbonates.

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

Compound	Trend
Hydroxides M(OH)₂	Solubility increases Mg → Ba
Sulfates MSO₄	Solubility decreases Mg → Ba (BaSO₄ insoluble — basis of sulfate test)

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.

MgCO₃ → MgO + CO₂ | 2Mg(NO₃)₂ → 2MgO + 4NO₂ + O₂

2Ba(NO₃)₂ → 2Ba(NO₂)₂ + O₂ (only Ba — nitrite formed, not oxide)

11Group 17 (Halogens)

Physical trends, oxidising/reducing ability, displacement reactions, and disproportionation of chlorine.

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⁻.

Cl₂ + 2KBr → 2KCl + Br₂ (orange)

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

Halide	AgNO₃ precipitate	Solubility
Cl⁻	White	Soluble in dilute NH₃
Br⁻	Cream	Soluble in conc. NH₃
I⁻	Pale yellow	Insoluble in NH₃

11.5 Disproportionation of Chlorine

Cold dilute NaOH (bleach)Cl₂ + 2NaOH → NaCl + NaOCl + H₂O

Hot conc. NaOH3Cl₂ + 6NaOH → 5NaCl + NaClO₃ + 3H₂O

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

12Nitrogen and Sulfur

N₂ unreactivity, NH₃ as a base, NO_x in the atmosphere, and acid rain chemistry.

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

Functional groups, IUPAC nomenclature, structural and stereoisomerism, hybridisation and curly-arrow mechanisms.

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

Hybrid	Geometry	Angle	Example
sp³	Tetrahedral	109.5°	CH₄
sp²	Trigonal planar	120°	C=C, C=O
sp	Linear	180°	C≡C, C≡N

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

Alkanes (free-radical substitution, cracking) and alkenes (electrophilic addition, Markovnikov, addition polymerisation).

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

Reagent / conditions	Product
H₂ / Pt or Ni, heat	Alkane (hydrogenation)
Steam / H₃PO₄, 300 °C, 60 atm	Alcohol
HX(g)	Halogenoalkane (Markovnikov)
Br₂(aq)	Bromohydrin (decolourises — test for C=C)
Cold dilute KMnO₄ / H⁺	Diol (purple → colourless)
Hot conc. KMnO₄	Cleaves C=C → carbonyls / CO₂

Markovnikov's rule: H adds to the C with more H atoms — the major product proceeds via the more stable secondary/tertiary carbocation (positive inductive effect of alkyl groups).

14.4 Addition Polymerisation

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

15Halogen Compounds

Reactions of halogenoalkanes (SN1, SN2, elimination) and relative reactivity by C–X bond strength.

15.1 Key Reactions

Reagent / conditions	Product	Type
NaOH(aq) / heat	Alcohol	Nucleophilic substitution
KCN in ethanol / heat	Nitrile (chain +1 C)	Nucleophilic substitution
NH₃ in ethanol / heat + pressure	Primary amine	Nucleophilic substitution
NaOH in ethanol / heat	Alkene	Elimination
AgNO₃ in ethanol / warm	Identification by AgX colour	—

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.

Common error: citing electronegativity to predict reactivity. The dominant factor is bond strength: weaker C–X means easier C–X breakage.

16Hydroxy Compounds (Alcohols)

Classification, oxidation to carbonyls/acids, dehydration to alkenes, esterification, and the iodoform test.

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⁺.

Alcohol	Distil immediately	Reflux
Primary	Aldehyde	Carboxylic acid
Secondary	Ketone	Ketone (no further oxidation)
Tertiary	No reaction — stays orange	No reaction

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

Aldehydes and ketones: nucleophilic addition with HCN, reduction, and distinguishing tests.

17.1 Common Reactions

Reagent	Aldehyde	Ketone
NaBH₄ or LiAlH₄ / dry ether	1° alcohol	2° alcohol
HCN / KCN(cat.)	2-hydroxynitrile	2-hydroxynitrile
2,4-DNPH	Orange ppt	Orange ppt
Tollens' reagent [Ag(NH₃)₂]⁺	Silver mirror	No reaction
Fehling's solution	Brick-red Cu₂O ppt	No reaction
I₂ / NaOH	Yellow CHI₃ ppt (only if CH₃CHO)	Yellow CHI₃ ppt (only if methyl ketone)

17.2 Mechanism of HCN Addition (Nucleophilic Addition)

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

If carbonyl is asymmetric (e.g. CH₃CHO), product has a chiral C → racemic mixture (50:50) because CN⁻ attacks both faces with equal probability.

18Carboxylic Acids and Esters

Acid–base properties, esterification, ester hydrolysis (acid and alkali), and reduction to alcohols.

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

Conditions	Products	Reversible?
Dilute acid / heat	Acid + alcohol	Reversible
Dilute alkali / heat (saponification)	Carboxylate salt + alcohol	Irreversible

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)

Primary amines, nitriles and hydroxynitriles: production and hydrolysis.

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)

Addition polymers from alkenes, repeat unit identification, and environmental impact.

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)

Designing multi-step routes between functional groups using AS reagents.

21.1 Strategy

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

21.2 Key Conversions

From	To	Reagents / conditions
Alkene	Halogenoalkane	HBr(g)
Halogenoalkane	Alcohol	NaOH(aq), reflux
Alcohol	Aldehyde	K₂Cr₂O₇/H⁺, distil
Aldehyde	Carboxylic acid	K₂Cr₂O₇/H⁺, reflux
Acid + alcohol	Ester	conc. H₂SO₄ cat.
Halogenoalkane	Nitrile (chain +1C)	KCN in ethanol, reflux
Nitrile	Carboxylic acid	H₂SO₄(aq), reflux

22Analytical Techniques (AS — IR & MS)

Identifying functional groups by IR absorption and deducing molecular structure from mass spectra.

22.1 Infrared Spectroscopy

Bond	Group	Wavenumber / cm⁻¹
C–O	alcohol, ester	1040–1300
C=C	alkene, arene	1500–1680
C=O	aldehyde, ketone, acid	1670–1740
C=O	ester	1710–1750
C≡N	nitrile	2200–2250
C–H	alkane	2850–2950
O–H (broad)	carboxylic acid	2500–3000
N–H	amine, amide	3300–3500
O–H	alcohol	3200–3600

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).

Isotope of Cl: M:M+2 ≈ 3:1; isotope of Br: M:M+2 ≈ 1:1.

These notes are AI-assisted study material. Always cross-check against the official 9701 syllabus or your teacher before relying on them in an exam.

23Chemical Energetics (A2)

Lattice energy and the Born–Haber cycle, hydration enthalpies, entropy and Gibbs free energy.

23.1 Lattice Energy

ΔH_latt⦵: the enthalpy change when 1 mole of an ionic solid is formed from its gaseous ions under standard conditions. Always exothermic.

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_f⦵(MX) = ΔH_at(M) + IE(M) + ΔH_at(X) + EA(X) + ΔH_latt⦵(MX)

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_sol = −ΔH_latt + ΣΔH_hyd

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

23.4 Entropy

Entropy S: a measure of the number of ways energy and particles can be arranged in a system.

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

23.5 Gibbs Free Energy

ΔG⦵ = ΔH⦵ − TΔS⦵

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)

Electrolysis quantitative work, standard electrode potentials, cell EMF, and the Nernst equation.

24.1 Electrolysis Calculations

Q = It | n(e⁻) = Q / F | F = 96 500 C mol⁻¹

mass = (M × It) / (n × F)

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

E⦵_cell = E⦵(more positive) − E⦵(more negative)

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

ΔG⦵ = −nFE⦵_cell

24.4 Nernst Equation

E = E⦵ + (0.059/z) lg([oxidised]/[reduced])

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

Common pitfall: writing ΔG⦵ = +nFE⦵_cell. The sign is negative — a positive cell EMF gives a negative ΔG (feasible).

25Equilibria (A2)

pH calculations, K_a, buffers, solubility product K_sp and partition coefficient.

25.1 pH

pH = −lg[H⁺] | [H⁺] = 10^−pH | K_w = [H⁺][OH⁻] = 1.00 × 10⁻¹⁴ at 298 K

25.2 Weak Acids

K_a = [H⁺][A⁻]/[HA] | [H⁺] ≈ √(K_a × [HA]) | pK_a = −lg K_a

25.3 Buffers

A solution that resists pH change when small amounts of acid or base are added. Acidic buffer: weak acid + its salt (e.g. CH₃COOH + CH₃COONa).

Henderson–HasselbalchpH = pK_a + lg ([A⁻]/[HA])

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

K_sp = [Mⁿ⁺]^a[X^m−]^b (units depend on stoichiometry)

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

25.5 Partition Coefficient K_pc

K_pc = [solute in solvent 1] / [solute in solvent 2]

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

26Reaction Kinetics (A2)

Rate equations, order, half-life, rate-determining step and types of catalysis.

26.1 Rate Equation

rate = k[A]^m[B]ⁿ

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

26.2 Orders

Order	[A] vs t	t½	Useful plot
Zero	Straight (gradient = −k)	Depends on [A]₀	[A] vs t
First	Curve, constant t½	0.693/k	ln[A] vs t (gradient = −k)
Second	Curve, t½ ↑ as time passes	1/(k[A]₀)	1/[A] vs t (gradient = +k)

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)

Thermal stability and solubility trends explained quantitatively using ionic radius and charge density.

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

d-block: variable oxidation states, coloured complexes, ligands, d-orbital splitting, and stability constants.

28.1 Definition

A transition element is a d-block element that forms at least one stable ion with an incomplete d sub-shell. Excludes Sc (only Sc³⁺ = d⁰) and Zn (Zn²⁺ = d¹⁰).

28.2 Characteristic Properties

Variable oxidation states — 3d and 4s electrons are similar in energy.
Catalytic behaviour — vacant d orbitals; ability to change oxidation state.
Form complex ions — vacant d orbitals accept lone pairs from ligands.
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⁻.

Example

[Cu(H₂O)₆]²⁺ pale blue + excess NH₃ → [Cu(NH₃)₄(H₂O)₂]²⁺ deep blue (ligand exchange).

28.5 Key Redox Reactions

2Cu²⁺ + 4I⁻ → 2CuI(s, white) + I₂ (disproportionation evidence; I₂ turns starch blue)

MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 4H₂O + 5Fe³⁺ (purple → colourless titration)

2MnO₄⁻ + 16H⁺ + 5C₂O₄²⁻ → 2Mn²⁺ + 8H₂O + 10CO₂ (slow at start, autocatalysed by Mn²⁺)

28.6 Stability Constant K_stab

The equilibrium constant for the formation of a complex from its constituent ions/molecules in solution. Larger K_stab → more stable complex (favours ligand exchange).

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

Additional functional groups, aromatic chemistry preview, and reactivity considerations.

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)

Delocalised structure, electrophilic substitution (nitration, halogenation, Friedel–Crafts), directing effects, and side-chain oxidation.

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

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

30.3 Key Reactions

Reaction	Reagents / conditions	Electrophile
Nitration	conc. HNO₃ + conc. H₂SO₄, 25–60 °C	NO₂⁺
Halogenation	Cl₂ / AlCl₃ (or Br₂ / FeBr₃)	Cl⁺ / Br⁺
Friedel–Crafts alkylation	R–Cl / AlCl₃	R⁺
Friedel–Crafts acylation	RCOCl / AlCl₃	RCO⁺

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

Unreactivity compared with halogenoalkanes — explained by overlap of halogen lone pair with the ring π system.

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

Acidity, electrophilic substitution at the ring, and reaction with diazonium salts.

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.

C₆H₅OH + 3Br₂ → C₆H₂Br₃OH + 3HBr

32.4 Azo Coupling

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

33Acyl Chlorides

Very reactive carbonyl derivatives — addition–elimination mechanism with O- and N-nucleophiles.

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

Nucleophile	Product
H₂O	Carboxylic acid + HCl
ROH	Ester + HCl
ArOH (phenol)	Aryl ester + HCl
NH₃ (excess)	Primary amide + NH₄Cl
R–NH₂	N-substituted amide + HCl

33.3 Mechanism (Addition–Elimination)

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

34Nitrogen Compounds (A2)

Aliphatic and aromatic amines, amides, diazonium salts, and amino acids.

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⁻).

Isoelectric point: the pH at which the net charge on the molecule is zero.

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

Polyesters and polyamides — small-molecule loss, biodegradability and structure prediction.

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)

Designing multi-step routes for aromatic and chiral molecules; analysing synthesis of pharmaceuticals.

36.1 Strategy

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

Chromatographic separation and ¹H / ¹³C NMR for structure determination.

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.

R_f = distance moved by spot / distance moved by solvent front

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:

Environment	δ / ppm
sp³ alkyl	0–50
C–O / C–N	50–90
sp² C=C / arene	110–160
C=O (carbonyl)	190–220

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.

Environment	δ / ppm
R–CH₃, –CH₂–	0.9–1.7
CH adjacent to C=O	2.2–3.0
CH adjacent to O / halogen	3.2–4.0
=CH (alkene)	4.5–6.0
Ar–H	6.0–9.0
RCHO (aldehyde)	9.3–10.5
RCOOH	9.0–13.0 (broad)

Common error: applying n+1 to H on the same carbon. Only count H on adjacent carbons.

📋Data, Constants and Exam Information

Reference cards for the AS & A Level examinations.

Fundamental Constants

Quantity	Symbol	Value
Avogadro constant	L	6.022 × 10²³ mol⁻¹
Molar gas constant	R	8.31 J K⁻¹ mol⁻¹
Faraday constant	F	9.65 × 10⁴ C mol⁻¹
Ionic product of water	K_w	1.00 × 10⁻¹⁴ mol² dm⁻⁶ (298 K)
Molar volume of gas	V_m	24.0 dm³ mol⁻¹ (room cond.); 22.4 dm³ mol⁻¹ (s.t.p.)
Specific heat capacity of water	c	4.18 J g⁻¹ K⁻¹
Electronic charge	e	−1.60 × 10⁻¹⁹ C

Paper Structure

Paper	Type	Time	Marks	% of AS	% of A
1	Multiple Choice	1 h 15	40	31%	15.5%
2	AS Structured	1 h 15	60	46%	23%
3	Advanced Practical Skills	2 h	40	23%	11.5%
4	A Level Structured	2 h	100	—	38.5%
5	Planning, Analysis & Evaluation	1 h 15	30	—	11.5%

Command Words

Word	Requirement
State	Fact only — no explanation
Define	Precise meaning
Describe	Key features, step-by-step
Explain	Give reasons using "because"
Calculate	Full working, with units
Deduce	Conclude using given evidence
Suggest	Apply knowledge to unfamiliar context
Compare	Similarities AND differences

Paper 5 — Linearisation Cheatsheet

Form	Plot	Gradient	Intercept
y = mx + c	y vs x	m	c
y = axⁿ	lg y vs lg x	n	lg a
y = a e^kx	ln y vs x	k	ln a

Cation Tests (Paper 3 — Qualitative Analysis)

Cation	NaOH(aq)	NH₃(aq)
Al³⁺	White ppt, soluble in excess	White ppt, insoluble in excess
Ca²⁺	White ppt (concentrated)	No ppt
Cu²⁺	Pale blue ppt, insoluble	Pale blue ppt, soluble → deep blue
Fe²⁺	Green ppt, brown in air	Same
Fe³⁺	Red-brown ppt, insoluble	Same
Mg²⁺	White ppt, insoluble	White ppt, insoluble
Zn²⁺	White ppt, soluble in excess	White ppt, soluble in excess
NH₄⁺	NH₃ on warming	—

Anion & Gas Tests

Species	Test
CO₃²⁻	CO₂ on adding dilute acid
Cl⁻ / Br⁻ / I⁻	AgNO₃: white / cream / pale yellow ppt
SO₄²⁻	BaCl₂ → white ppt, insoluble in HCl
NH₃ (gas)	Damp red litmus turns blue
CO₂	Limewater milky
H₂	Squeaky pop with lit splint
O₂	Relights glowing splint

Exam Tips

For enthalpy definitions, always state "1 mole … standard conditions … standard states" — examiners look for these specific phrases.

For ionic equations, only include species that actually change; remove spectator ions.

When asked to "explain" a trend, state the property first, then the cause (nuclear charge, radius, shielding), then the effect.

For Le Chatelier, never write that a catalyst shifts the equilibrium position — it does not.

In rate equation questions, never deduce orders from stoichiometry — always use experimental data.

9701 AS Level Chemistry

1Atomic Structure

1.1 Sub-atomic Particles

1.2 Electron Sub-shells and Orbitals

1.3 Shapes of s and p Orbitals

1.4 Ionisation Energy

Worked example

2Atoms, Molecules and Stoichiometry

2.1 Key Definitions

2.2 Mole Calculations

2.3 Yield and Limiting Reagent

2.4 Ionic Equations

Example

2.5 Common Ions to Memorise

3Chemical Bonding

3.1 Electronegativity

3.2 Ionic Bonding

3.3 Covalent and Coordinate Bonding

3.4 Metallic Bonding

3.5 VSEPR Shapes

3.6 Intermolecular Forces

4States of Matter

4.1 The Ideal Gas

4.2 Lattice Types

4.3 Allotropes of Carbon

5Chemical Energetics (AS)

5.1 Enthalpy Change ΔH

5.2 Standard Enthalpy Definitions

5.3 Calorimetry

5.4 Bond Energies

5.5 Hess's Law

6Electrochemistry (AS)

6.1 Oxidation Number Rules

6.2 Oxidation and Reduction

6.3 Balancing Redox Equations

Worked example

7Equilibria (AS)

7.1 Dynamic Equilibrium

7.2 Le Chatelier's Principle

7.3 Kc and Kp

7.4 Brønsted–Lowry Acids and Bases

7.5 Industrial Processes

8Reaction Kinetics (AS)

8.1 Collision Theory

8.2 Boltzmann Distribution

9Chemical Periodicity (Period 3)

9.1 Physical Trends Na → Ar

9.2 Elements with Oxygen and Water

9.3 Period 3 Chlorides with Water

10Group 2

10.1 Reactivity Trends Mg → Ba

10.2 Solubility Trends

10.3 Thermal Stability

11Group 17 (Halogens)

11.1 Physical Properties Cl₂ → I₂

11.2 Halogens as Oxidising Agents

11.3 Halide Ions as Reducing Agents

11.4 AgNO₃ Test for Halides

11.5 Disproportionation of Chlorine

12Nitrogen and Sulfur

12.1 Inertness of N₂

12.2 Ammonia as a Base

12.3 NOx in the Atmosphere

12.4 Acid Rain

13Introduction to AS Organic Chemistry

13.1 Functional Groups (AS)

13.2 Nomenclature

13.3 Isomerism

13.4 Hybridisation

13.5 Reaction Types

13.6 Curly Arrows

14Hydrocarbons

14.1 Alkanes

14.2 Free-Radical Substitution with X₂/UV

14.3 Alkenes — Electrophilic Addition

14.4 Addition Polymerisation

15Halogen Compounds

15.1 Key Reactions

15.2 SN1 vs SN2

15.3 C–X Reactivity

7.3 K_c and K_p

12.3 NO_x in the Atmosphere

25.5 Partition Coefficient K_pc

27.2 Solubility Explained by ΔH_sol

28.6 Stability Constant K_stab