Advanced Level Chemistry – Quick Reference Notes

Unit 1 – Atomic Structure

• Matter = anything with mass + occupies space → ~118 known elements → ~100 found naturally.
• Historical ideas
– Empedocles: 4 elements.
– Democritus: atomos (indivisible).
– Dalton (1808) ⮕ Golf-ball model; 4 postulates (atoms, identity, conservation, fixed ratios).

Discovery of sub-atomic particles

• Cathode rays → electrons (J.J. Thomson 1897) → $e/m =1.76\times10^{8}\,\text{C g}^{-1}$.
• Millikan oil-drop (1909) → $q<em>e =1.602\times10^{-19}\,\text{C}$; $m</em>e =9.11\times10^{-28}\,\text{g}$.
• Canal (anode) rays → protons, $m<em>p =1.673\times10^{-24}\,\text{g}$. • Rutherford gold-foil → nuclear model; nucleus tiny, +vely charged. • Chadwick (1932) → neutron, $m</em>n =1.675\times10^{-24}\,\text{g}$, $q=0$.

Atomic number Z, mass number A

• $Z=p=e$ (neutral atom).
• $A=Z+n$.
• Isotopes: same Z, different A.
• Unified amu $1\,\text{u}= \dfrac1{12}M(^{12}\text C)=1.66054\times10^{-24}\,\text g$.
• Average atomic mass $\bar{M}=\sum (m<em>i\,x</em>i)$.

Electromagnetic radiation

• Wave parameters: $c=\lambda\nu$; $E=h\nu$ where $h=6.626\times10^{-34}\,\text{J s}$.
• Einstein photon; de Broglie matter wave $\lambda=\dfrac{h}{mv}$.

Hydrogen spectrum & Bohr model

• Electrons occupy orbits $n=1,2\dots$; energy quantised → line spectra (Lyman, Balmer, Paschen…).

Quantum numbers

• n (principal), l (0 → n–1; s p d f), $m<em>l=-l..+l$, $m</em>s=\pm\tfrac12$.
• Subshell degeneracy; total orbitals per shell $n^2$.

Electron configurations

• Aufbau order (1s,2s,2p,3s,3p,4s,3d…).
• Pauli Exclusion: max 2e⁻ per orbital opposite spin.
• Hund: degenerate orbitals singly with parallel spins first.
• Condensed form e.g. $[\text{Ne}]3s^1$.
• Cr, Cu anomalies: $[Ar]3d^54s^1$, $[Ar]3d^{10}4s^1$.

Periodic Table building

• Groups 1–18 (IUPAC).
• Blocks: s (2 columns), p (6), d (10), f (14).

Periodic trends (s & p blocks)

• Atomic radii ↓ across (Z_eff ↑), ↑ down (n ↑).
• Ionisation energy ↑ across, ↓ down; big jumps after core e⁻ removal.
• Electron gain (affinity) most exothermic for halogens; less favourable down group.
• Electronegativity (Pauling) ↑ across, ↓ down.

Unit 2 – Structure & Bonding

Types of bonds

• Ionic: electrostatic M⁺ … X⁻, lattice E.
• Covalent: shared pair; octet rule; $\sigma,\;\pi$ bonds.
• Dative (coordinate): both electrons donated e.g. \ce{NH3->BF3}.
• Metallic: cations in e⁻ sea; strength ↑ with charge density.

Lewis structures

• Procedure: count valence e⁻, place pairs, satisfy octet/duet, formal charges $FC=V-(L+B)$ minimise.
• Resonance ↔ equivalent contributors; hybrid ↓E.

VSEPR geometries (AXmEn)

• 2 e-pairs: linear 180° (BeCl2, CO2).
• 3: trig-planar 120° (BF3); AX2E bent <120° (SO2).
• 4: tetrahedral 109.5° (CH4); AX3E trig pyramidal (NH3); AX2E2 bent 104.5° (H2O).
• 5: trig bipyramidal (PCl5); see-saw (SCl4); T-shape (ICl3); linear (XeF2).
• 6: octahedral (SF6); square pyramidal (XeOF4); square planar (XeF4).

Hybridisation

• sp (linear), sp2 (trig planar), sp3 (tetrahedral).
• Multiple bonds: one $\sigma$ + one/two $\pi$.

Inter-molecular forces

• Ion-dipole, H-bond (FON), dipole-dipole, London (dispersion), dipole-induced, ion-induced.
• Boiling/melting relate to IMF & molar mass (n-pentane vs neopentane).

Unit 3 – Chemical Calculations

Oxidation number rules

• Free element 0; Group1 +1; Group2 +2; F −1; H +1 (−1 in hydrides); O −2 (−1 in peroxides, −½ in superoxides, +2 in OF2).
• Sum in molecule 0, in ion = charge.

Redox concepts

• Oxidation = loss e⁻ (ON ↑), Reduction = gain e⁻ (ON ↓).
• Balancing: inspection; oxidation-number; half-reaction (acid/base media).

Formulae & composition

• Empirical from %: convert to mol, divide by smallest, scale ↓.
• Molecular: $\text{n}=\dfrac{M<em>r}{M</em>{emp}}$.
• Mass %, ppm, ppb; $\text{ppm}=\dfrac{\text{mass solute}}{\text{mass solution}}10^{6}$.
• Molality $m=\dfrac{n<em>{solute}}{kg</em>{solvent}}$.
• Molarity $c=\dfrac{n}{V(\text{dm}^3)}$.

Stoichiometry

• Limiting reagent concept.
• Gravimetric & titrimetric calculations.

Solution prep

• Weigh solid → dissolve → volumetric flask.
• Dilution $M<em>1V</em>1=M<em>2V</em>2$.

Unit 6 – Chemistry of s, p & d Blocks

Group 1 (Alkali)

• Valence $ns^1$; low density, soft, MP ↓ down group.
• Reactivity ↑ down; with \ce{H2O}: $2M+2H2O→2MOH+H2$ (Li slow, Cs explosive).
• O₂: $4M+O2→2M2O$; K/Rb/Cs form superoxides $MO2$.
• Salts: all soluble; LiF, Li2CO3 sparingly.
• Flame colours: Li red, Na yellow, K lilac.

Group 2 (Alkaline earth)

• Valence $ns^2$; harder, MP high.
• Water: Be none; Mg hot water; Ca→Ba vigorous.
• Thermal stability of carbonates ↑ down; nitrates give $M(NO<em>3)</em>2 → MO+NO2+O2$.
• Flame: Ca orange-red, Sr crimson, Ba apple-green.

p-Block highlights

• Group13: B non-metal, Al amphoteric; $Al<em>2O</em>3$ amphoteric; $AlCl<em>3$ dimeric $Al</em>2Cl<em>6$. • Group14: C allotropes (diamond sp3, graphite sp2, $\pi$ deloc), fullerenes; $CO$ triple bond; $CO</em>2$ linear.
• Group15: N₂ inert (triple 942 kJ mol⁻¹); variable ox states −3→+5; $NH<em>3$ basic ligand; $HNO3$ strong oxidiser. • Group16: O₂/ O₃ oxidants; S allotropes $S</em>8$; acid–base nature of oxides shift basic→acidic across period; hydrides acidity ↑ down (HF weak due H-bond).
• Group17: Halogens oxidising strength F2>Cl2>Br2>I2; HX acid strength ↑ down; $ClO^-$ series (-ite, –ate); $KMnO<em>4$ tests (purple→colourless). • Group18: Noble gases; Xe forms $XeF</em>2, XeF<em>4, XeO</em>3,$ etc.

d-Block (Transition)

• Definition: incomplete d in atom or common ion; Sc (d¹) – Zn (d¹⁰).
• Properties: variable ox states (+2 prevalent), coloured complexes, catalytic, magnetic, form alloys.
• Electronic anomalies Cr (3d⁵4s¹), Cu (3d¹⁰4s¹).
• Complex nomenclature: ligand names (ammine, aqua, chlorido); e.g. $[Fe(CN)<em>6]^{3-}$ = hexacyanidoferrate(III). • Colour factors: metal, ox state, ligand field. • Qualitative tests: – \ce{Fe^{2+}} + \ce{[Fe(CN)6]^{3-}} → Turnbull blue.
– \ce{Fe^{3+}} + \ce{SCN^-} → blood-red $[Fe(SCN)]^{2+}$.
– \ce{Cr^{3+}} + \ce{H2O2/OH^-} → yellow $CrO_4^{2-}$.

Key Equations (selection)

• $c=\lambda\nu$; $E=h\nu$; $\lambda=\dfrac{h}{mv}$
• Atomic mass $\bar{M}=\sum m<em>ix</em>i$.
• Molality $m=\dfrac{n}{kg}$; Molarity $M=\dfrac{n}{V<em>{dm^3}}$. • Ideal dilution $M</em>1V<em>1=M</em>2V_2$.
• Gibbs $\Delta G=\Delta H-T\Delta S$.

Ethical / Practical Notes

• H₂S, CO, Cl₂ highly toxic → lab ventilation.
• KMnO₄ stains; concentrated $H<em>2SO</em>4$ dehydrating agent.
• Transition-metal waste can be environmentally harmful; proper disposal.