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.

Shapes and orientations of s, p and d atomic orbitals — Atomic-orbital shapes and orientations. Use the p-orbital row to distinguish p_x, p_y and p_z; click the diagram to enlarge.

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.

1.5 Atomic and Ionic Radius Trends

Atomic radius decreases across a period (same shell, increasing nuclear charge) and increases down a group (new shell, more shielding). A cation is smaller than its atom; an anion is larger. A beam of protons is deflected towards the negative plate, electrons more strongly the other way (larger charge/mass), neutrons not at all.

1.6 Writing Ionisation-Energy Equations

First IE: X(g) → X⁺(g) + e⁻. Second IE starts from the ion: X⁺(g) → X²⁺(g) + e⁻. Use successive IE jumps to deduce the group and the electronic configuration.

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)

Titration apparatus showing initial and final burette readings and the colour change at the end point — Read the burette at eye level before and after titration; the titre is final reading minus initial reading, and the end point is the first permanent indicator colour change.

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⁺.

2.6 Definitions You Must State

Unified atomic mass unit: 1/12 of the mass of one ¹²C atom.
A_r, relative isotopic mass, M_r and relative formula mass are all measured on that scale.
One mole contains L = 6.02×10²³ specified particles.
Empirical formula is the simplest whole-number ratio; molecular formula is a whole-number multiple of it.
Hydrated salts contain water of crystallisation; anhydrous salts do not.

Mole calculations must cover reacting masses, gas volumes (24.0 dm³ mol⁻¹ at r.t.p. if given), solution concentrations, percentage yield and limiting reagent. Always use state symbols in ionic equations and omit spectators.

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.

Sideways overlap of p orbitals forming a pi bond — A π bond forms by sideways overlap of adjacent p orbitals, leaving electron density above and below the internuclear axis.

Period 3 and below can expand the octet (use d orbitals), giving SO₃, 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.

Positive metal ions surrounded by a sea of delocalised outer-shell electrons — Metallic bonding is non-directional: positive ions are held together by attraction to mobile delocalised electrons, which also carry charge through the lattice.

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.

An instantaneous dipole inducing a dipole in a neighbouring atom and producing attraction — Random electron movement creates an instantaneous dipole; it induces an opposite dipole nearby, so the two particles attract temporarily.

Hydrogen bonding between three polar water molecules — Hydrogen bonds form between a δ⁺ hydrogen on one molecule and a lone pair on an electronegative atom of another molecule.

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.

3.7 σ and π Bonds; Hybridisation

A σ bond is end-on overlap of orbitals on the internuclear axis. A π bond is sideways overlap of p orbitals. H₂ is σ only; C₂H₄ is one σ + one π; N₂ and HCN have a triple bond (one σ + two π).

sp³ (tetrahedral, 109.5°) in CH₄ and C₂H₆; sp² (120°) in C₂H₄; sp (180°) in HCN and C₂H₂.

3.8 Expanded Octet and Dative Bonds

Period 3 can expand the octet: SO₂, PCl₅, SF₆. A dative (coordinate) bond is a shared pair provided by one atom: NH₃ + HCl → NH₄⁺; Al₂Cl₆ is two AlCl₃ units linked by dative bonds from Cl to Al.

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

Three representations of alternating sodium and chloride ions in a giant ionic lattice — A giant ionic lattice extends in three dimensions. Each ion is surrounded by oppositely charged ions, producing strong electrostatic attraction throughout the crystal.

Regular layers sliding in a pure metal compared with distorted layers in an alloy — Equal-sized atoms in a pure metal form layers that slide readily. Different-sized atoms in an alloy disrupt the layers, making sliding harder and the material stronger.

4.3 Allotropes of Carbon

Structures of graphite, diamond and fullerene — Graphite, diamond and fullerene contain only carbon but differ in bonding and structure, producing very different physical properties.

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 (C...): simple molecular sphere; held by id-id forces; low mp.

4.4 Ideal and Real Gases

pV = nRT

Real gases deviate at high pressure (finite molecular volume) and low temperature (intermolecular attractions). They behave most ideally at low p and high T.

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 and 101 kPa, with substances in their standard states.

In both profiles the peak is the activation barrier. Products lie below reactants for an exothermic reaction and above them for an endothermic reaction.

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".

Insulated polystyrene cup calorimeter with lid and thermometer — A lidded polystyrene cup reduces heat exchange with the surroundings. Measure the largest temperature change and use the mass of solution in *q = mcΔT*.

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.

General Hess cycle linking reactants, products and elements — General Hess cycle: choose a common intermediate, keep arrow directions consistent, and apply the algebraic sign of each route.

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

5.6 Calorimetry Equation and Standard Conditions

q = mcΔT then ΔH = ±q/n

Standard enthalpy changes refer to 101 kPa and a stated temperature, usually 298 K, with standard states. Exothermic ΔH is negative. Average bond energy is the mean enthalpy to break one mole of that bond in the gaseous state.

6Electrochemistry (AS)

Redox language, oxidation numbers, oxidising and reducing agents, and balancing half-equations.

6.1 Redox Definitions

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

The two always occur together. Older criteria (gain/loss of oxygen or hydrogen) agree with the oxidation-number test when all four can be applied; the oxidation-number test is the one examiners expect.

Disproportionation: the same element is oxidised and reduced in one reaction (Cl₂ + cold dilute NaOH). Comproportionation is the reverse: two different oxidation states of one element meet at an intermediate state.

6.2 Oxidation Number Rules

Uncombined element (including O₂, Cl₂, S₈): 0.
Group 1 in compounds: +1. Group 2: +2. F: always −1.
O: usually −2; peroxides −1; OF₂ +2.
H: usually +1; metal hydrides (NaH) −1.
Monoatomic ion: charge on the ion.
Sum in a neutral compound is 0; sum in a polyatomic ion equals the ion charge.
The more electronegative atom takes the negative number.

Use Stock names: iron(II) chloride is FeCl₂; chlorate(I) is ClO⁻; nitrate(V) is NO₃⁻; manganate(VII) is MnO₄⁻. Everyday “nitrate” means nitrate(V).

6.3 Oxidising and Reducing Agents

An oxidising agent accepts electrons and is itself reduced (MnO₄⁻/H⁺, Cr₂O₇²⁻/H⁺, Cl₂, conc. H₂SO₄, O₂). A reducing agent donates electrons and is itself oxidised (I⁻, Fe²⁺, H₂, C, reactive metals, SO₂).

H₂O₂ is in the intermediate O(−1) state: it oxidises when oxygen goes to −2 (extra water formed) and reduces when O₂ is produced.

Worked example

MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O.
Mn +7 → +2 (reduced; MnO₄⁻ is the oxidant). Fe +2 → +3 (oxidised; Fe²⁺ is the reductant). The titration is self-indicating: first permanent pale pink is the end-point.

6.4 Balancing Redox Equations

Use this order in acid: balance the redox element; add H₂O for O; add H⁺ for H; add e⁻ for charge; multiply halves so electrons cancel; never leave free electrons in the overall equation.

MnO₄⁻ → Mn²⁺ needs 4H₂O on the right, 8H⁺ on the left, then 5e⁻ on the left.
Fe²⁺ → Fe³⁺ + e⁻. Multiply by 5 and add.
Check atoms and total charge on both sides.

In alkali, add OH⁻ to both sides to convert leftover H⁺ into water. Concentrated H₂SO₄ oxidises Br⁻ to Br₂ and is reduced to SO₂ — so it cannot be used to dry HBr.

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.

Concentration-time graphs showing the same equilibrium reached from reactants or from products — The final equilibrium composition is the same whichever side is used initially. Concentrations become constant, but both forward and reverse reactions continue at equal rates.

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: pK_a 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.

7.6 Writing K_c and K_p

K_c uses equilibrium concentrations; K_p uses partial pressures (mole fraction × total pressure). Pure solids and liquids do not appear. K changes only with temperature. A catalyst does not change K — it shortens the time to reach the same position.

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.

Reaction flask connected to a gas syringe for measuring gas volume against time — For a gas-producing reaction, record gas-syringe volume at regular times. The initial gradient of the volume-time graph gives the initial rate.

8.2 Boltzmann Distribution

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

Boltzmann distributions at two temperatures with activation energy marked — Raising temperature lowers and broadens the distribution peak; the area to the right of E_a increases substantially.

↑ 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.

Catalysed and uncatalysed reaction pathways with different activation energies but the same enthalpy change — A catalyst lowers activation energy by providing another pathway. It does not change reactant or product energies, so ΔH and the equilibrium position are unchanged.

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.

8.3 Catalysts

A catalyst provides an alternative route with lower E_a. On a Boltzmann diagram, more molecules then exceed the new E_a. Draw the catalysed profile as a lower peak (sometimes two peaks if there is an intermediate). The catalyst is regenerated.

Homogeneous — same phase as the reactants (e.g. Fe²⁺/Fe³⁺ in I⁻/S₂O₈²⁻).
Heterogeneous — different phase, usually a solid surface (Haber Fe; Contact V₂O₅).

Rate is change of concentration per unit time. From a concentration–time graph, gradient at a point is the instantaneous rate; use initial rate for comparing conditions.

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 the 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

9.4 Period 3 Oxides and Chlorides

	Na	Mg	Al	Si	P	S
Oxide	Na₂O	MgO	Al₂O₃	SiO₂	P₄O₁₀	SO₂/SO₃
Bonding	ionic	ionic	ionic/covalent	giant covalent	simple covalent	simple covalent
Oxide + water	strong alkali	weak alkali	insoluble	insoluble	strong acid	strong acid
Chloride + water	neutral pH~7	slightly acidic	acidic (fumes)	acidic	acidic	—

Al₂O₃ and Al(OH)₃ are amphoteric: they dissolve in both acid and NaOH. Na and Mg react with water; later elements do not. Oxidation numbers of the oxides/chlorides match the number of valence electrons used.

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 because atomic radius increases, so the 2 outer electrons are 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)

10.4 Reactions to Recall

Mg → Ba react with O₂, water (Mg needs steam) and dilute HCl/H₂SO₄ to give H₂. Oxides, hydroxides and carbonates all react with dilute acid to give a salt + water (+ CO₂ from carbonates). Hydroxide solubility increases down the group; sulfate solubility decreases. Thermal stability of carbonates and nitrates increases down the group (larger cation, less polarising).

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 decreases down the group (stronger id-id forces as electron count increases).
Bond energy Cl-Cl > Br-Br > I-I (longer bond — weaker overlap).

Halogen samples showing pale green chlorine, orange bromine and purple iodine colours — Halogen colours become deeper down Group 17: chlorine is pale green, bromine orange-brown and iodine gives a purple vapour.

11.2 Halogens as Oxidising Agents

Strength: Cl₂ > Br₂ > I₂. Cl₂ displaces Br₂ and I₂; Br₂ displaces only I₂.

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.

11.6 Hydrogen Halides and Concentrated H₂SO₄

Thermal stability of HX decreases I > Br > Cl because the H–X bond weakens. Concentrated H₂SO₄ on solid halides: Cl⁻ gives HCl only; Br⁻ is oxidised to Br₂ (brown, SO₂); I⁻ is oxidised to I₂ (purple vapour, H₂S / S). This is the halide-as-reducing-agent trend.

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).

Dot-and-cross diagram of ammonium showing nitrogen donating both electrons in the coordinate bond — In NH₄⁺, nitrogen donates its lone pair to H⁺. After formation, all four N-H bonds are equivalent even though one began as a coordinate bond.

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.

12.5 Haber and Contact Processes

Haber: N₂ + 3H₂ ⇌ 2NH₃, Fe catalyst, ~450 °C, ~200 atm. Contact: 2SO₂ + O₂ ⇌ 2SO₃, V₂O₅, ~450 °C, ~1–2 atm; SO₃ is dissolved in conc. H₂SO₄ then diluted (never add SO₃ straight to water).

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.

13.7 Stereoisomerism at AS

Cis–trans / E–Z needs a C=C (no rotation) and two different groups on each carbon. Optical isomerism needs a chiral carbon (four different groups); a pair of non-superimposable mirror images rotates plane-polarised light equally in opposite directions.

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

Initiation, propagation and termination steps in chlorination of methane — The propagation steps consume one radical and create another, so the chain continues. Termination removes radicals when two collide.

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

Summary of hydrogenation, hydration, hydrogen-halide addition and halogen addition across a carbon-carbon double bond — The C=C π bond opens during addition. The atoms from the reagent attach across the two carbon atoms, producing a saturated product.

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).

14.5 Making Alkenes

Elimination of HX from a halogenoalkane: ethanolic NaOH, heat.
Dehydration of an alcohol: heated Al₂O₃ or conc. H₂SO₄.
Cracking of a longer alkane over Al₂O₃.

14.6 Electrophilic Addition — Reagents and Markovnikov

H₂ / Ni or Pt, heat → alkane.
Steam / H₃PO₄ → alcohol.
HX(g), r.t. → halogenoalkane. The H adds to the carbon already having more H (more stable 2°/3° carbocation; alkyl groups are electron-releasing).
X₂ → dihalogenoalkane. Br₂(aq) decolorised is the C=C test.
Cold dilute acidified KMnO₄ → diol (purple to colourless).
Hot concentrated acidified KMnO₄ cleaves the C=C; the fragments tell you where the double bond was.

14.7 Alkanes — Combustion and Cracking

Alkanes are unreactive to polar reagents (strong, non-polar C–H). Incomplete combustion gives CO and C. Catalytic converters remove CO, NO_x and unburnt hydrocarbons. Cracking heavier fractions gives more useful shorter alkanes and alkenes.

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.

15.4 Nucleophilic Substitution — Reagents

NaOH(aq), heat → alcohol.
KCN in ethanol, heat → nitrile (one extra carbon).
NH₃ in ethanol, heat under pressure → amine.
AgNO₃(aq) in ethanol: rate of AgX ppt identifies the halogen (C–I fastest, weakest bond).

Ethanolic NaOH, heat → elimination to an alkene. Primary halogenoalkanes favour S_N2; tertiary favour S_N1.

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

Ethanol and acidified potassium dichromate heated gently while ethanal is distilled off — Distil the aldehyde as it forms to prevent further oxidation. Reflux instead keeps it with the oxidising agent, producing the carboxylic acid.

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.

16.5 Making Alcohols and Further Reactions

Routes: steam + alkene / H₃PO₄; NaOH(aq) + halogenoalkane; NaBH₄ or LiAlH₄ on a carbonyl; LiAlH₄ on a carboxylic acid; hydrolysis of an ester.

Na(s) → sodium alkoxide + H₂. Alcohols are weaker acids than water.
PCl₅, SOCl₂ or HX → halogenoalkane (steamy fumes with PCl₅ test the OH).
Conc. H₂SO₄ or Al₂O₃, heat → alkene (dehydration).
Carboxylic acid + conc. H₂SO₄ → ester.
Acidified K₂Cr₂O₇: distill a 1° alcohol to the aldehyde; reflux to the acid. 2° → ketone. 3° no reaction (stays orange).

17Carbonyl Compounds

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

17.1 Common Reactions

Polar carbonyl group with a partially positive carbon and partially negative oxygen — Oxygen withdraws electron density from C=O, leaving carbon δ⁺. Nucleophiles therefore attack the carbonyl carbon.

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.

17.3 Distinguishing Tests

2,4-DNPH: orange/yellow ppt with any carbonyl (aldehyde or ketone).
Tollens' (ammoniacal AgNO₃, warm): silver mirror with aldehydes only.
Fehling's / Benedict's: brick-red Cu₂O with aliphatic aldehydes.
Alkaline I₂: yellow CHI₃ ppt if a CH₃CO− or CH₃CH(OH)− group is present.

Reduction with NaBH₄ or LiAlH₄ gives the corresponding alcohol. HCN / KCN, heat gives a hydroxynitrile (nucleophilic addition — extra carbon).

18Carboxylic Acids and Esters

Naming, hydrogen bonding, acid reactions, esterification and hydrolysis.

18.1 Structure and Physical Properties

The carboxyl group is −COOH. Methanoic, ethanoic, propanoic acids are the first three. Two molecules dimerise through two hydrogen bonds, so boiling points are higher than alcohols of similar M_r. Short-chain acids mix with water; longer chains do not.

18.2 Reactions of Carboxylic Acids

Na or Mg → salt + H₂.
NaOH or Na₂CO₃/NaHCO₃ → salt + H₂O (+ CO₂ with carbonate). Effervescence with carbonate distinguishes them from phenol and alcohols.
Alcohol + conc. H₂SO₄ / heat → ester + H₂O (reversible).
LiAlH₄ in dry ether, then acid → primary alcohol. NaBH₄ does not reduce carboxylic acids.
PCl₅ or SOCl₂ → acyl chloride (steamy HCl fumes with PCl₅).

18.3 Why They Are Acidic

Stronger than alcohols and phenol because the carboxylate ion is resonance-stabilised: the negative charge is shared equally over both oxygen atoms. Still weak acids: only partially dissociated in water.

18.4 Esters

Named as alkyl alkanoates (CH₃COOCH₂CH₃ is ethyl ethanoate). Pleasant smells; used as solvents and flavourings.

Hydrolysis	Conditions	Products	Reversible?
Acid	dilute H₂SO₄, heat	acid + alcohol	Yes — excess water needed
Alkali (saponification)	dilute NaOH, heat	carboxylate salt + alcohol	No — salt does not re-esterify

18.5 Making Carboxylic Acids

Reflux a 1° alcohol or an aldehyde with acidified K₂Cr₂O₇ / KMnO₄.
Hydrolyse a nitrile (dilute acid, or alkali then acidify).
Hydrolyse an ester, then acidify if alkali was used.

19Nitrogen Compounds (AS)

Primary amines, nitriles and hydroxynitriles: production, basicity, hydrolysis and reduction.

19.1 Primary Amines

Amines are derivatives of ammonia in which one or more H atoms are replaced by alkyl or aryl groups. A primary amine has one alkyl/aryl group attached to the nitrogen atom: R-NH₂.

Basicity: amines act as bases because the lone pair on N can accept a proton (H⁺) from water or acid.

Order of base strength: ethylamine > ammonia > phenylamine

Ethylamine: the ethyl group is electron-donating (+I effect), pushing electron density towards N — lone pair more available — stronger base.
Ammonia: three H atoms on N; no electron-donating alkyl group.
Phenylamine: the lone pair on N is delocalised into the benzene ring — less available to accept H⁺ — weakest base.

19.2 Production of Primary Amines

Method	Reagents / conditions	Equation
From halogenoalkane	Excess NH₃ in ethanol, heat under pressure	R-X + NH₃ → R-NH₂ + HX
From nitrile	LiAlH₄ in dry ether; or H₂ / Ni catalyst	R-CN + 4[H] → R-CH₂NH₂
From amide	LiAlH₄ in dry ether	R-CONH₂ + 4[H] → R-CH₂NH₂ + H₂O

Two-step formation of a primary amine from a halogenoalkane and ammonia — Ammonia first attacks the halogenoalkane to form an alkylammonium ion; a second ammonia molecule removes H⁺, giving the primary amine.

Excess ammonia is essential when making amines from halogenoalkanes to prevent further substitution: secondary amine (R₂NH), tertiary amine (R₃N) and quaternary ammonium salt (R₄N⁺X⁻) can all form.

19.3 Nitriles

Nitriles contain the -C≡N functional group. They are useful for chain extension (adding one carbon atom to the chain).

Formation: R-X + KCN in ethanol, reflux → R-CN + KX (nucleophilic substitution; chain lengthens by 1 C).
Hydrolysis: reflux with dilute HCl or H₂SO₄ → carboxylic acid + NH₄⁺ salt.

CH₃CH₂CN + 2H₂O + HCl → CH₃CH₂COOH + NH₄Cl
Reduction: LiAlH₄ / dry ether or H₂ / Ni → primary amine.

CH₃CN + 4[H] → CH₃CH₂NH₂

19.4 Hydroxynitriles

Carbonyl compounds (aldehydes and ketones) react with HCN (in the presence of a base catalyst such as KCN or NaCN) to form 2-hydroxynitriles (cyanohydrins).

CH₃CHO + HCN → CH₃CH(OH)CNethanal → 2-hydroxypropanenitrile

Hydroxynitriles are useful intermediates:

Hydrolysis of the -CN group → α-hydroxy carboxylic acid (e.g. CH₃CH(OH)CN + 2H₂O + H⁺ → CH₃CH(OH)COOH + NH₄⁺).
Reduction of the -CN group → amino alcohol (e.g. CH₃CH(OH)CN + 4[H] → CH₃CH(OH)CH₂NH₂).

If the carbonyl carbon is part of an asymmetric molecule, the product contains a chiral centre — racemic mixture formed because CN⁻ attacks both faces of the planar C=O with equal probability.

20Polymerisation (AS)

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

20.1 Addition Polymerisation

Addition polymerisation involves the joining together of many alkene monomer units. The C=C double bond opens up and the monomers link together to form long chains. No small molecules are lost.

n CH₂=CH₂ → [−CH₂−CH₂−]_npoly(ethene)

Repeat unit: the smallest repeating section of the polymer chain. To draw a repeat unit from a monomer: open the C=C and put brackets around the two-carbon fragment.

Ethene monomers becoming the bracketed repeat unit of polyethene — Open the monomer C=C to form two single bonds that continue through the brackets; retain the substituents on their original carbon atoms.

Monomer	Repeat unit	Polymer name	Common use
ethene	[−CH₂−CH₂−]	poly(ethene)	plastic bags, bottles
chloroethene	[−CH₂−CHCl−]	poly(chloroethene) (PVC)	pipes, window frames
propene	[−CH₂−CH(CH₃)−]	poly(propene)	crates, rope
phenylethene	[−CH₂−CH(C₆H₅)−]	poly(phenylethene) (polystyrene)	packing, insulation
tetrafluoroethene	[−CF₂−CF₂−]	poly(tetrafluoroethene) (PTFE)	non-stick coatings

20.2 Deducing the Monomer from a Repeat Unit

Work backwards: replace the single C-C bond in the repeat unit with a C=C double bond, and remove the brackets. Any side groups stay attached to the same carbon.

Example: repeat unit [−CH₂−CHCl−] → monomer is CH₂=CHCl (chloroethene).

20.3 Environmental Issues

Non-biodegradable: C-C backbone is chemically inert; bacteria and fungi cannot break it down. Polymers persist in landfill for hundreds of years.
Toxic combustion products: PVC releases HCl; polyurethanes release HCN; burning polystyrene releases soot and CO.
Solutions: recycling (mechanical: melting and remoulding; chemical: cracking back to monomers), biodegradable alternatives (e.g. polylactic acid from corn starch), and energy recovery (incineration for electricity generation).

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) or HCl(g)
Alkene	Alcohol	Steam / H₃PO₄, 300 °C, 60 atm
Alkene	Diol	Cold dilute KMnO₄ / H⁺
Halogenoalkane	Alcohol	NaOH(aq), reflux
Halogenoalkane	Alkene	NaOH in ethanol, heat
Halogenoalkane	Nitrile (chain +1C)	KCN in ethanol, reflux
Halogenoalkane	Primary amine	NH₃ in ethanol, heat, pressure
Alcohol	Aldehyde	K₂Cr₂O₇/H⁺, distil
Alcohol	Carboxylic acid	K₂Cr₂O₇/H⁺, reflux
Alcohol	Halogenoalkane	PCl₅ or SOCl₂ or KBr + conc. H₂SO₄
Alcohol	Alkene	conc. H₂SO₄, 170 °C or Al₂O₃, hot
Alcohol	Ester	RCOOH + conc. H₂SO₄ cat., warm
Aldehyde	Primary alcohol	NaBH₄ or LiAlH₄ / dry ether
Aldehyde	Carboxylic acid	K₂Cr₂O₇/H⁺, reflux
Ketone	Secondary alcohol	NaBH₄ or LiAlH₄ / dry ether
Carboxylic acid	Primary alcohol	LiAlH₄ / dry ether
Carboxylic acid	Ester	ROH + conc. H₂SO₄ cat., warm
Carboxylic acid	Carboxylate salt	NaOH, Na₂CO₃, or reactive metal
Nitrile	Carboxylic acid	Dilute HCl or H₂SO₄, reflux
Nitrile	Primary amine	LiAlH₄ / dry ether or H₂ / Ni

When planning a synthesis, always check the carbon count. Converting a halogenoalkane to a nitrile adds one carbon to the chain (R-X → R-CN → R-CH₂NH₂ or R-COOH).

21.3 Multi-step Routes

Identify every functional group, then list the type of reaction and the exact reagent/condition for each step. Watch for extra carbons (KCN) and for oxidation that can overshoot (distil vs reflux). Possible by-products: further substitution of amines, over-oxidation of aldehydes, elimination competing with substitution.

22Analytical Techniques (AS — IR & MS)

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

22.1 Infrared Spectroscopy

Infrared spectrum of ethanol with a broad alcohol O-H absorption — IR spectrum of ethanol: identify the broad alcohol O-H absorption around 3200-3600 cm⁻¹ before using the fingerprint region.

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

Mass spectrometer showing vaporisation, ionisation, acceleration, deflection and detection — Mass-spectrometer sequence: vaporise, ionise, accelerate, deflect in a magnetic field, then detect ions according to m/z.

M⁺ peak — molecular mass.
[M+1]⁺ from ¹³C: n_C = 100 × I(M+1) / (1.1 × I(M)).
[M+2]⁺: M:M+2 ≈ 3:1 for one Cl (³⁵Cl:³⁷Cl), and ≈ 1:1 for one Br (⁷⁹Br:⁸¹Br).
Fragmentation: loss of 15 (CH₃), 17 (OH), 29 (CHO or C₂H₅), 31 (OCH₃), 43 (C₃H₇ or CH₃CO), 45 (COOH or OC₂H₅).

Mass spectra of propanone and propanal showing different fragment peak patterns — Compounds with the same molecular formula can share the same molecular-ion mass but give different fragment patterns; those patterns act as structural fingerprints.

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

22.3 Using the Data-Booklet Peaks

IR: match strong absorptions to O−H (broad, acids very broad), C=O (~1700 cm⁻¹), C−O, C≡N, C−H. Absence of a peak can be as useful as its presence.

MS: M⁺ peak gives M_r. The [M+1]⁺ peak height estimates the number of carbons. Simple fragments (15 = CH₃, 29 = C₂H₅ / CHO, 77 = C₆H₅) identify pieces. A_r from isotope abundances is a weighted mean.

ASRevision Guide Exam Boosters

Condensed additions from the AS part of the local CIE 9701 complete revision guide: practical wording, calculation discipline, reaction conditions and common examiner traps.

AS Paper 1/2 high-yield checks

Area	What to write in exam answers	Common trap
Atomic structure	Use proton number to define the element; use mass number minus proton number for neutrons. For ionisation energy definitions, include gaseous atoms, one mole and one electron.	For transition-metal ions, remove 4s electrons before 3d electrons.
Stoichiometry	Convert volumes to dm³ for concentration calculations and temperatures to K for pV = nRT. Use the balanced equation mole ratio before deciding the limiting reagent.	Using the reagent in excess as the limiting reagent because it has the larger starting mass.
Bonding and structure	Link property to structure and bonding: type of particles, forces between them, and whether mobile charged particles are present.	Saying "strong intermolecular forces" for giant ionic, metallic or covalent structures.
Energetics	For Hess cycles, keep arrows consistent. For calorimetry, quote heat loss, incomplete combustion and evaporation as likely reasons experimental values are less exothermic.	Forgetting that bond breaking is endothermic and bond making is exothermic.
Equilibrium and rates	Only temperature changes K_c or K_p. Catalysts increase both forward and reverse rates equally. Temperature increases rate mainly by greatly increasing the fraction with E ≥ E_a.	Claiming a catalyst changes equilibrium yield or that rate orders come from equation coefficients.

AS organic reaction conditions

Conversion	Reagents and conditions	Observation / exam note
Alkane to halogenoalkane	Cl₂ or Br₂, ultraviolet light	Free-radical substitution; mixture of products likely.
Alkene to alcohol	Steam, H₃PO₄ catalyst, high temperature and pressure	Hydration; addition across C=C.
Alkene test	Br₂(aq), room temperature	Orange/brown to colourless; electrophilic addition.
Halogenoalkane to alcohol	NaOH(aq), heat under reflux	Nucleophilic substitution.
Halogenoalkane to alkene	NaOH or KOH in ethanol, heat	Elimination; forms H₂O and halide ion.
Primary alcohol to aldehyde	Acidified K₂Cr₂O₇, distil	Distillation prevents further oxidation.
Primary alcohol to carboxylic acid	Acidified K₂Cr₂O₇, reflux	Use excess oxidising agent; orange to green.
Carbonyl to hydroxynitrile	HCN generated from KCN/NaCN and dilute acid	Planar C=O gives a racemic mixture if a chiral centre forms.
Nitrile to amine	LiAlH₄ in dry ether, then water	Adds one carbon before reduction when made from halogenoalkane + CN^-.

AS practical wording

Titration reliability: rinse burette with the solution it will contain, remove funnel before titrating, read the meniscus at eye level, and repeat until concordant titres are obtained.
Calorimetry improvements: use a lid, insulation, a draught shield and a more precise thermometer; extrapolate cooling curves when required.
Qualitative analysis: describe precipitate colour, solubility in excess reagent, gas test and final inference. Do not infer an ion from colour alone.
Uncertainty: percentage uncertainty = absolute uncertainty / measured value x 100. For repeated measurements, use half the range as an estimate when appropriate.

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

Born-Haber cycle for the formation of lithium fluoride — Born-Haber cycle for LiF. Check the physical state of every species and the sign of each enthalpy change before solving for lattice enthalpy.

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 +).

Worked example: Born-Haber for NaCl

Given: ΔH_f(NaCl) = −411 kJ mol⁻¹; ΔH_at(Na) = +107; IE(Na) = +496; ΔH_at(Cl) = +122; EA(Cl) = −349. Calculate ΔH_latt.

ΔH_f = ΔH_at(Na) + IE(Na) + ΔH_at(Cl) + EA(Cl) + ΔH_latt
−411 = 107 + 496 + 122 + (−349) + ΔH_latt
ΔH_latt = −411 − 376 = −787 kJ mol⁻¹

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)

Electrolytic cell connected to a direct-current supply with positive anode and negative cathode — In an electrolytic cell, the external supply makes the anode positive and the cathode negative. Oxidation occurs at the anode and reduction at the cathode.

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).

Chlorine half-cell connected by a salt bridge to a standard hydrogen electrode — Connect the unknown half-cell to the standard hydrogen electrode. Platinum provides an inert surface, while the salt bridge completes the circuit without mixing the solutions directly.

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])

How to read the equation at 298 K:

More oxidised form → [oxidised]/[reduced] gets larger → E becomes more positive → reduction is favoured.
More reduced form → [oxidised]/[reduced] gets smaller → E becomes more negative.

Why a positive cell EMF is feasibleE^°_cell > 0 ⇒ −nFE^°_cell < 0 ⇒ ΔG^° < 0 ⇒ feasible

Remember the minus sign: ΔG^° = −nFE^°_cell, not +nFE^°_cell. Both n and F are always positive.

24.5 Predicting Electrolysis Products

Molten binary salts: metal at the cathode, non-metal at the anode. In aqueous solution, water can be discharged: H₂ at the cathode if the metal is more reactive than hydrogen; O₂ at the anode unless concentrated halide is present (then Cl₂/Br₂/I₂). Dilute vs concentrated NaCl is a standard contrast.

F = Le Q = It moles of e⁻ = Q/F

Avogadro's constant can be found by measuring the mass of metal deposited and the charge passed.

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

Hydrogen chloride and water forming hydronium and chloride with conjugate pairs marked — Conjugate acid-base pairs differ by exactly one H⁺: HCl/Cl⁻ and H₂O/H₃O⁺.

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.

25.6 Calculations

pH = −lg[H⁺]; K_w = [H⁺][OH⁻] = 1.00×10⁻¹⁴ at 298 K. Strong acid: [H⁺] = concentration. Strong alkali: [OH⁻] = concentration, then [H⁺] = K_w/[OH⁻]. Weak acid: [H⁺] = √(K_a × c) if α is small. A buffer is a weak acid + its salt (or weak base + salt); added H⁺ is removed by A⁻, added OH⁻ by HA. K_sp = product of aqueous ion concentrations raised to stoichiometric powers — solids omitted.

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	t_1/2	Useful plot
Zero	Straight (gradient = −k)	Depends on [A]₀	[A] vs t
First	Curve, constant t_1/2	0.693/k	ln[A] vs t (gradient = −k)
Second	Curve, t_1/2 increases as time passes	1/(k[A]₀)	1/[A] vs t (gradient = +k)

Rate against concentration curves for zero-order, first-order and second-order reactions — On a rate-versus-concentration graph: zero order is horizontal, first order is linear through the origin, and second order curves upward.

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 Arrhenius Equation

k = Ae^−E_a/RT | ln k = ln A − E_a / RT

Plot ln k vs 1/T gives gradient = −E_a/R, intercept = ln A. Larger E_a → smaller k at a given T; higher T → larger k.

The exponential term dominates: a small increase in T produces a large increase in k because the fraction of molecules with E ≥ E_a rises exponentially.

26.5 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⁻.

Diffusion, adsorption, surface reaction, desorption and diffusion in iron-catalysed ammonia formation — Heterogeneous catalysis proceeds by diffusion to the surface, adsorption and bond weakening, reaction on active sites, desorption, then diffusion away.

26.6 First-Order Half-Life and Finding Orders

For a first-order reaction t_1/2 is constant and k = ln 2 / t_1/2. Orders come from initial-rate data (how rate changes when one [ ] is doubled) or from a concentration–time shape (zero: straight; first: constant half-life). The rate-determining step's molecularity must match the rate equation; intermediates formed in fast steps before the r.d.s. can appear in the rate equation after substitution.

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.

Compound	Mg	Ca	Sr	Ba
Carbonate decomposition T	~540 °C	~900 °C	~1290 °C	~1360 °C
Nitrate decomposition T	~450 °C	~560 °C	~570 °C	~590 °C

Group 2 carbonates: MCO₃ → MO + CO₂. Nitrates (except Ba): 2M(NO₃)₂ → 2MO + 4NO₂ + O₂. Ba(NO₃)₂ → Ba(NO₂)₂ + O₂ (only Ba gives nitrite not oxide).

The higher the charge density of the cation, the greater the polarising power — Mg²⁺ has the highest charge density in Group 2, so its compounds decompose at the lowest temperatures.

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.

Key exam comparison

BaSO₄ is insoluble — basis of the sulfate ion test (BaCl₂ + H⁺ gives white ppt). MgSO₄ and CaSO₄ are soluble.
Mg(OH)₂ is sparingly soluble (milk of magnesia); Ba(OH)₂ is much more soluble → stronger alkaline solution.

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

Shapes and orientations of the five d orbitals — The five d orbitals have different orientations relative to approaching ligands; this determines how their energies split in a complex.

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⁻ (ethylenediaminetetraacetate) — forms very stable 1:1 complexes via the chelate effect.
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⁻.

Octahedral d orbital splitting and visible-light absorption — Octahedral splitting creates two energy levels. Absorption of visible light promotes an electron across ΔE; the complementary colour is observed.

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 (redox 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⁺.

28.8 Catalysis and Variable Oxidation States

Transition elements have an incomplete d subshell in the atom or a common ion. Variable oxidation states let them catalyse redox (e.g. Fe²⁺/Fe³⁺, MnO₂ for H₂O₂, V₂O₅ in Contact). Ligands are Lewis bases; denticity is the number of donor atoms (en is bidentate, EDTA⁴⁻ hexadentate). Colour arises when a d–d transition absorbs visible light — the complementary colour is seen. A ligand change or oxidation-state change changes Δ and therefore the colour.

29Introduction to A2 Organic Chemistry

Additional functional groups, aromatic chemistry, and reactivity considerations that extend AS knowledge.

29.1 New Functional Groups (A2)

At A2 you need to recognise and name compounds containing these additional functional groups:

Functional group	General formula	Example
Arene (benzene ring)	C₆H₅-	benzene, methylbenzene
Halogenoarene	C₆H₅-X	chlorobenzene, bromobenzene
Phenol	C₆H₅-OH	phenol
Acyl chloride	RCOCl	ethanoyl chloride
Amide	RCONH₂	ethanamide
Amino acid	H₂N-CHR-COOH	glycine, alanine
Azo (diazo)	-N=N-	azobenzene dyes

29.2 Key Reactivity Patterns

Benzene vs alkenes: Benzene undergoes electrophilic substitution (aromaticity preserved); alkenes undergo electrophilic addition. The delocalised π system in benzene is more stable than isolated C=C bonds.
Phenol reactivity: The -OH group donates electron density into the ring via resonance — phenol is more reactive than benzene towards electrophiles (e.g. reacts with Br₂(aq) without catalyst).
Acyl chlorides: Most reactive carboxylic acid derivative due to strongly polar C=O and excellent leaving group (Cl⁻). Reacts with H₂O, alcohols, phenols and amines via addition-elimination.
Amides: Much less reactive than acyl chlorides; the N lone pair is delocalised into the C=O bond, reducing electrophilicity of the carbonyl carbon.

29.3 Bonding and Hybridisation (A2 Review)

Review of AS hybridisation with A2 extensions:

Hybrid	Geometry	Angle	Example
sp³	Tetrahedral	109.5°	CH₄, C-OH in alcohols
sp2	Trigonal planar	120°	C=C, C=O, benzene ring carbons
sp	Linear	180°	C≡C, C≡N

In benzene, each C is sp2 hybridised. The remaining p orbital on each C overlaps side-on with neighbours to form a delocalised π system of 6 electrons above and below the ring plane.

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.

Benzene sigma framework and overlapping p orbitals forming delocalised pi clouds above and below the ring — Each sp² carbon retains one p orbital. Sideways overlap around the entire ring creates delocalised π electron density above and below the plane.

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⁺ converts R-C₆H₅ to benzoic acid (regardless of side-chain length, provided it has at least one benzylic H).

30.6 Benzene Reactions — Conditions

Nitration: conc. HNO₃ + conc. H₂SO₄, ~55 °C. Electrophile NO₂⁺.
Halogenation: Cl₂ / AlCl₃ or FeCl₃, r.t.
Friedel–Crafts alkylation / acylation: RCl or RCOCl / AlCl₃.
Side-chain oxidation of any alkylbenzene with alkaline KMnO₄ then acid → benzoic acid (needs a benzylic H).

Benzene does not decolorise Br₂(aq) and does not react with H₂ without a catalyst — evidence against a Kekulé cyclohexatriene.

31Halogenoarenes

Unreactivity of aryl halides compared with alkyl halides — explained by overlap of a halogen lone pair with the ring π system.

31.1 Structure and Bonding

In halogenoarenes (e.g. chlorobenzene, bromobenzene), the halogen atom is directly bonded to a carbon of the benzene ring. The C-X bond is shorter and stronger than in halogenoalkanes.

31.2 Unreactivity towards Nucleophiles

Chlorobenzene does not react with aqueous NaOH (unlike chloroethane, which readily hydrolyses). Two reasons explain this:

Lone pair delocalisation: A lone pair on the halogen atom overlaps with the ring π system, giving partial double-bond character in the C-X bond. This shortens and strengthens the C-X bond, making it harder to break.
Electron repulsion: The delocalised π electron cloud of the ring repels approaching nucleophiles, hindering attack at the carbon.

Chlorine lone-pair overlap with the benzene pi system producing partial double-bond character — A chlorine lone pair overlaps with the ring π system. The resulting partial double-bond character shortens and strengthens C-Cl, so nucleophilic substitution is difficult.

31.3 Comparison with Halogenoalkanes

Feature	Halogenoalkane (e.g. C₂H₅Cl)	Halogenoarene (e.g. C₆H₅Cl)
C-X bond length	Longer	Shorter (partial double bond character)
C-X bond strength	Weaker	Stronger
Reaction with NaOH(aq)	Readily hydrolyses	No reaction
Mechanism	SN1 or SN2	Not applicable under normal conditions

31.4 Side-chain Halogenation

When alkylbenzenes are treated with Cl₂ or Br₂ in the presence of UV light, substitution occurs in the side-chain (not the ring) by a free-radical mechanism. This produces halogenoalkylbenzenes (e.g. C₆H₅CH₂Cl from methylbenzene).

C₆H₅CH₃ + Cl₂ → C₆H₅CH₂Cl + HCluv light, side-chain substitution

This is a useful distinction: halogenation in the ring requires AlCl₃/FeBr₃ catalyst (electrophilic substitution); halogenation in the side-chain requires UV light (free-radical substitution).

32Phenol

Acidity compared with ethanol and carboxylic acids, and activated-ring substitution.

32.1 Acidity

Phenol is more acidic than ethanol because the phenoxide ion is stabilised by delocalisation of the negative charge into the ring. It is less acidic than a carboxylic acid, where the charge is shared by two oxygen atoms. Phenol dissolves in NaOH(aq) but not in Na₂CO₃(aq) — a standard test against carboxylic acids.

32.2 Reactions of the −OH Group

Na(s) → sodium phenoxide + H₂.
NaOH(aq) → sodium phenoxide + H₂O.
No reaction with carbonates.

32.3 Electrophilic Substitution

The OH group activates the ring and directs 2- and 4-. Phenol decolorises Br₂(aq) at room temperature with no catalyst, giving a white ppt of 2,4,6-tribromophenol and HBr. Dilute HNO₃ gives a mixture of 2- and 4-nitrophenol.

C₆H₅OH + 3Br₂ → C₆H₂Br₃OH + 3HBr

32.4 Azo Coupling

Phenoxide + benzenediazonium chloride in alkali below 10 °C → orange/yellow azo dye. The −N=N− link is the chromophore. Used in dyes and as a qualitative test for a phenol.

33Acyl Chlorides

Why they are the most reactive carboxylic-acid derivatives, and the addition–elimination mechanism.

33.1 Reactivity

The carbonyl carbon is strongly δ+ and Cl⁻ is a good leaving group. Reactivity order: acyl chloride > acid anhydride > ester > amide. Acyl chlorides fume in moist air (HCl).

33.2 Key Reactions (room temperature, often vigorous)

Nucleophile	Organic product	Inorganic product
H₂O	carboxylic acid	HCl
ROH	ester	HCl
phenol / NaOH	aryl ester	HCl / NaCl
NH₃ (excess)	primary amide	NH₄Cl
RNH₂	N-substituted amide	RNH₃Cl

33.3 Addition–Elimination

Nucleophile attacks the δ+ carbonyl carbon; the C=O π bond breaks onto oxygen (tetrahedral intermediate).
The C=O bond reforms and Cl⁻ leaves.
A proton is lost if the nucleophile was neutral.

Compared with carboxylic acids, no heat or conc. acid is needed, and the reactions go to completion because HCl leaves.

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.

34.6 Amino Acids and Amides in More Detail

α-Amino acids exist as zwitterions in the solid and at the isoelectric pH. In acid they are cations; in alkali, anions. They join by condensation to peptides. Amides are neutral (the lone pair is delocalised onto the C=O). Hydrolysis of an amide (acid or alkali, heat) recovers the carboxylic acid (or salt) and the amine (or ammonium salt).

35Condensation Polymerisation

Polyesters, polyamides, repeating units and why they can be hydrolysed.

35.1 Polyesters

A diol + a dicarboxylic acid (or dioyl chloride) lose H₂O (or HCl) at each link. Terylene / PET is benzene-1,4-dicarboxylic acid + ethane-1,2-diol. The repeating unit contains the ester −COO− group.

35.2 Polyamides

A diamine + a dicarboxylic acid (or dioyl chloride) give a polyamide. Nylon-6,6 is hexane-1,6-diamine + hexanedioic acid. Kevlar uses 1,4-aromatic monomers and hydrogen-bonds between chains, giving high tensile strength. Proteins are natural polyamides (peptide bonds).

35.3 Predicting the Repeating Unit

Remove H from each −OH or −NH₂ and OH from each −COOH (or Cl from an acyl chloride) and join the remainders. For a single amino acid or hydroxy-acid, the monomer already contains both functional groups.

35.4 Biodegradability

Ester and amide links can be hydrolysed by acid, alkali or enzymes, so polyesters and polyamides are potentially biodegradable. Addition polymers have a C−C backbone that is not easily hydrolysed.

36Organic Synthesis (A2)

Designing multi-step routes for aromatic and chiral molecules; retrosynthetic analysis and pharmaceutical synthesis.

36.1 Synthetic Strategy

Good synthesis planning requires careful consideration of:

Functional group interconversions: match target functional groups with known transformations from the reagent list.
Regioselectivity: in aromatic chemistry, directing effects determine where new groups attach (e.g. nitration of methylbenzene — 2-nitromethylbenzene and 4-nitromethylbenzene).
Stereoselectivity: consider chirality when creating or modifying chiral centres. Markovnikov addition, SN2 inversion, and HCN addition to carbonyls all create defined stereochemistry.
Order of steps: install reactive groups at the appropriate stage. For example, nitration of methylbenzene is done before oxidising the side-chain to COOH, because -COOH is a 3-director (meta-directing) while -CH₃ is a 2,4-director.
Protecting groups: sometimes a reactive group must be protected before another reaction. For example, -NH₂ can be acetylated to -NHCOCH₃ before nitration, then hydrolysed back afterwards.

36.2 Key A2 Synthetic Routes

Target	Route	Key reagents
Benzoic acid from methylbenzene	Side-chain oxidation	KMnO₄ / OH⁻, reflux; then H⁺
Phenylamine from methylbenzene	CH₃ → COOH → NO₂ → NH₂	KMnO₄; conc. HNO₃/H₂SO₄; Sn/HCl
2,4,6-tribromophenol from phenol	Direct bromination (no catalyst needed)	Br₂(aq), room temp
Nylon-6,6 from adipic acid	Dicarboxylic acid + diamine	Heat, condensation
Aspirin from salicylic acid	Acylation of -OH	(CH₃CO)₂O or CH₃COCl

Reaction map connecting benzene, bromobenzene, methylbenzene, benzoic acid, nitrobenzene, phenylamine and phenol — Use the map as a route-planning model: work backwards from the target, label every arrow with reagent and conditions, and check that directing effects permit the required substitution position.

36.3 Chiral Synthesis and Pharmaceuticals

Many drugs are chiral, and enantiomers can have very different biological effects:

Thalidomide: one enantiomer treats morning sickness; the other causes birth defects. Racemic mixture was dangerous.
Ibuprofen: only the (S)-enantiomer is active; the (R)-enantiomer is inactive (but converts slowly in the body).

Asymmetric hydrogenation creating one chiral form of naproxen using a chiral catalyst — A chiral catalyst creates one enantiomer preferentially when a planar group is converted into a tetrahedral stereocentre, reducing the need to separate a racemic mixture.

Methods to obtain single enantiomers:

Chiral catalysis: use a chiral catalyst or enzyme to favour one enantiomer (e.g. asymmetric hydrogenation).
Resolution: convert a racemic mixture into diastereomers (e.g. by reacting with a single enantiomer of a chiral acid), separate by crystallisation, then regenerate the original enantiomers.
Chiral pool synthesis: start from a naturally occurring chiral compound (e.g. amino acids, sugars).

In exam questions, always state whether a chiral centre is created, destroyed, or preserved at each step. If a new chiral centre is formed from a planar intermediate (e.g. C=O + HCN), a racemic mixture is produced.

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

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

37.1 Thin-Layer Chromatography (TLC)

Chromatography chamber with baseline, solvent and reference spots labelled — Chromatography setup: the baseline and sample spots must start above the solvent level, and the solvent front is marked immediately after removal.

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)

Gas-liquid chromatography instrument showing injector, column, carrier gas and detector — GLC apparatus: an inert carrier gas moves the vaporised sample through a coated column to a detector and recorder.

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.

Gas chromatogram with separated peaks at different retention times — Each peak appears at a characteristic retention time. Peak position helps identify a component; integrated peak area is proportional to its 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

Carbon-13 NMR spectrum of propylbenzene with its distinct carbon environments indicated — Each chemically distinct carbon environment gives one ¹³C signal. Symmetry can make several carbon atoms equivalent and reduce the number of peaks.

37.4 1H NMR

Low resolution proton NMR spectrum with signals integrating to one, two and three hydrogens — Low-resolution ¹H NMR: chemical shift identifies the environment, while relative peak area gives the 1H:2H:3H proton ratio.

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.

37.5 Reading an NMR Spectrum

TMS is the zero reference. Chemical shift tells the environment (near O or a C=O is more downfield).
¹H integration (peak area) is the proton ratio.
Splitting: n+1 rule from non-equivalent neighbouring protons. A singlet has no such neighbours.
¹³C counts different carbon environments; it is not split by H in the usual decoupled spectrum.
TLC: R_f = distance spot / distance solvent. GLC: retention time identifies; peak area is amount.

...Data, Constants and Exam Information

Reference cards for the AS & A Level examinations.

A2 revision guide exam boosters

A2 area	Exam-useful detail	Common trap
Lattice energy and Born-Haber cycles	Define lattice energy with gaseous ions and one mole of ionic solid. Use atomisation, first/second ionisation energy, electron affinity, enthalpy of formation and lattice energy with arrows in the correct direction.	Using aqueous ions in a lattice-energy definition, or forgetting that the second electron affinity of oxygen is endothermic.
Entropy and Gibbs free energy	ΔG = ΔH - TΔS. Convert ΔS from J K^-1 mol^-1 to kJ K^-1 mol^-1 when ΔH is in kJ mol^-1. Feasible means ΔG < 0.	Mixing units or saying a feasible reaction must be fast; feasibility is thermodynamic, not kinetic.
Electrode potentials	More positive E^° values favour reduction. E^°_cell = E^°_right - E^°_left, or cathode minus anode.	Changing the sign of E^° when multiplying a half-equation; electrode potentials are not multiplied.
A2 kinetics	Orders come only from experimental data. Units of k depend on overall order. The rate-determining step must contain the species that appear in the rate equation, possibly after replacing intermediates.	Deducing rate order from stoichiometric coefficients in the overall equation.
Transition elements	Variable oxidation states, coloured complexes, ligand exchange and catalysis come from partially filled d orbitals and close 3d/4s energy levels.	Calling Zn a transition element in 9701 definitions; Zn forms Zn²⁺ with full 3d¹⁰.
NMR and chromatography	Use number of peaks for environments, integration for relative proton count, splitting for adjacent H atoms, and D₂O exchange for O-H/N-H. In chromatography, retention time identifies; peak area estimates amount.	Applying the n+1 rule to hydrogens on the same carbon instead of adjacent carbons.

A2 organic synthesis checklist

Work backwards: identify the target functional group, then choose the immediate precursor with a known reagent set.
Count carbons: CN^- substitution and Friedel-Crafts acylation can extend the carbon skeleton; many oxidations and reductions do not.
Control ring position: alkyl and hydroxy groups direct mainly 2,4; nitro, carboxyl and acyl groups direct mainly 3. The order of steps can decide the final isomer.
Protect reactive groups when needed: convert -NH₂ to an amide before harsh electrophilic substitution if the question requires controlled ring substitution.
Track chirality: nucleophilic addition to planar carbonyl compounds produces a racemic mixture when a new chiral centre forms.

Paper 5 planning and analysis

Skill	What to include
Independent variable	Name it, give a sensible range and at least five values where a graph is expected.
Dependent variable	State exactly what is measured, with units and measuring instrument.
Controlled variables	Choose variables that would genuinely affect the result: temperature, concentration, volume, surface area, mass, catalyst amount or timing method.
Risk assessment	Pair each hazard with a specific precaution, for example corrosive acid with eye protection and immediate washing of spills.
Graph work	Label axes with quantity and unit, choose a scale using over half the grid, draw a best-fit line/curve, and calculate gradient from a large triangle.
Reliability	Repeat, identify anomalies, calculate a mean only from concordant or justified readings, and improve precision with appropriate apparatus.

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%

Complete Calculation and Formula Reference

Area	Formula / relationship	Non-negotiable condition
Amount of substance	n = m/M; n = cV; c = n/V; n = V/V_m; pV = nRT	For concentration use V in dm³; for the ideal-gas equation use pressure in Pa, volume in m³ and temperature in K. Use the balanced equation before applying mole ratios.
Atomic and mass spectrometry	Relative atomic mass = Σ(isotopic mass × abundance)/Σabundance; m/z = mass number/charge	Use fractional abundance or percentages consistently. A doubly charged ion has half the m/z value for the same mass.
Energetics	q = mcΔT; ΔH = -q/n; ΔH = ΣΔH_f(products) - ΣΔH_f(reactants); ΔH = Σbond energies broken - Σbond energies formed	For calorimetry, explain the minus sign from heat gained by water versus heat change of reaction. State that mean bond enthalpies are estimates.
Entropy and feasibility	ΔG = ΔH - TΔS	Use K. Convert J K⁻¹ mol⁻¹ to kJ K⁻¹ mol⁻¹ if ΔH is in kJ mol⁻¹. Feasible at stated conditions means ΔG < 0, not necessarily fast.
Electrochemistry	Q = It; n(e⁻) = Q/F; E^°_cell = E^°_cathode - E^°_anode	Time must be in s. Do not multiply an electrode potential when multiplying a half-equation. A positive standard cell emf indicates thermodynamic feasibility under standard conditions.
Equilibria and acids	K_c = [products]^coefficients/[reactants]^coefficients; K_p = partial pressures of products/reactants; pH = -lg[H⁺]; K_w = [H⁺][OH⁻]; K_a = [H⁺][A⁻]/[HA]; K_sp = product of equilibrium ion concentrations	Omit pure solids and liquids from equilibrium expressions. Equilibrium concentrations must be used, not initial values; indices are stoichiometric coefficients.
Kinetics	rate = change in amount/concentration per time; rate = k[A]^m[B]ⁿ; lnk = lnA - E_a/RT	Find orders only from experimental data. The units of k depend on overall order. For Arrhenius work, use K and consistent energy units.
Analytical chemistry	R_f = distance moved by solute/distance moved by solvent front; % yield = actual/theoretical ×100; % atom economy = M_r desired product/ΣM_r products ×100	Measure both TLC distances from the baseline. Use formula masses and a balanced equation for atom economy; yield can never legitimately exceed 100%.

Calculation method: write the governing equation, insert values with units, convert units before using a calculator, retain guard figures, and finish with a chemically meaningful unit and conclusion.

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

1.5 Atomic and Ionic Radius Trends

1.6 Writing Ionisation-Energy Equations

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

2.6 Definitions You Must State

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

3.7 σ and π Bonds; Hybridisation

3.8 Expanded Octet and Dative Bonds

4States of Matter

4.1 The Ideal Gas

4.2 Lattice Types

4.3 Allotropes of Carbon

4.4 Ideal and Real Gases

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

5.6 Calorimetry Equation and Standard Conditions

6Electrochemistry (AS)

6.1 Redox Definitions

6.2 Oxidation Number Rules

6.3 Oxidising and Reducing Agents

Worked example

6.4 Balancing Redox Equations

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

7.6 Writing Kc and Kp

8Reaction Kinetics (AS)

8.1 Collision Theory

8.2 Boltzmann Distribution

8.3 Catalysts

9Chemical Periodicity (Period 3)

9.1 Physical Trends Na → Ar

9.2 Elements with Oxygen and Water

9.3 Period 3 Chlorides with Water

9.4 Period 3 Oxides and Chlorides

10Group 2

10.1 Reactivity Trends Mg → Ba

10.2 Solubility Trends

10.3 Thermal Stability

10.4 Reactions to Recall

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

11.6 Hydrogen Halides and Concentrated H2SO4

12Nitrogen and Sulfur

12.1 Inertness of N2

12.2 Ammonia as a Base

12.3 NOx in the Atmosphere

12.4 Acid Rain

12.5 Haber and Contact Processes

13Introduction to AS Organic Chemistry

13.1 Functional Groups (AS)

7.3 K_c and K_p

7.6 Writing K_c and K_p

11.1 Physical Properties Cl₂ → I₂

11.4 AgNO₃ Test for Halides

11.6 Hydrogen Halides and Concentrated H₂SO₄

12.1 Inertness of N₂

12.3 NO_x in the Atmosphere

14.2 Free-Radical Substitution with X₂/UV

24.2 Standard Electrode Potential E^°