Chemistry — Last-Minute Revision
30 must-know Chemistry points covering all 13 chapters. Read through this the night before or morning of your TGBIE 1st Year Chemistry exam.
Bohr model: radius rₙ = n²a₀/Z (increases with n, decreases with Z). Energy Eₙ = −13.6Z²/n² eV (negative = bound state).
Balmer series is visible; Lyman is UV. 'Hydrogen spectrum lines': Lyman→Balmer→Paschen→Brackett→Pfund (UV→vis→IR).
Anomalous electronic configs: Cr = [Ar]3d⁵4s¹ (half-filled d = extra stability); Cu = [Ar]3d¹⁰4s¹ (fully-filled d = extra stability).
Atomic radius DECREASES across a period; INCREASES down a group.
Ionisation enthalpy increases across a period (harder to remove electron from smaller, more nuclear-attracted atom).
VSEPR: NH₃ = trigonal pyramidal (1 lone pair); H₂O = bent (2 lone pairs); CH₄ = tetrahedral (0 lone pairs).
Bond order in MOT: O₂ bond order = 2 (double bond), has 2 unpaired e⁻ → paramagnetic. N₂ bond order = 3 → diamagnetic.
Hydrogen bond strength: F−H···F > O−H···O > N−H···N. High BP of H₂O is due to intermolecular H-bonding.
Ideal gas: PV = nRT. R = 8.314 J mol⁻¹K⁻¹ = 0.082 L·atm mol⁻¹K⁻¹.
Graham's law: heavier gas diffuses SLOWER. r ∝ 1/√M. H₂ diffuses ~4× faster than O₂.
Van der Waals: 'a' corrects for intermolecular attraction (lowers pressure); 'b' corrects for finite molecular volume.
Molarity changes with temperature (volume changes); molality does NOT change with temperature (mass of solvent is constant).
Redox: oxidation = loss of electrons = increase in oxidation number. Reduction = gain of electrons.
ΔH = ΔU + ΔnRT where Δn = moles of gaseous products − moles of gaseous reactants.
Hess's law: ΔH is independent of path. Add or subtract reactions with their ΔH values algebraically.
Spontaneity: ΔG < 0. ΔG = ΔH − TΔS. If ΔH < 0 and ΔS > 0 → always spontaneous (all T).
Kc: pure solids and liquids NOT included. Kp = Kc(RT)^Δn where Δn = moles of gaseous products − reactants.
Le Chatelier: increase P → equilibrium shifts toward fewer moles of gas. Increase T → shifts toward endothermic side.
pH = −log[H⁺]; pOH = −log[OH⁻]; pH + pOH = 14 at 25°C. pH 7 = neutral, <7 = acidic, >7 = basic.
Ksp = [A^m+]^a[B^n-]^b. Common ion effect: adding a common ion DECREASES solubility.
Alkali metals (Li, Na, K...): increase in reactivity down the group; Li anomalous (diagonal relationship with Mg).
Na + excess O₂ → Na₂O₂ (peroxide). Li + O₂ → Li₂O (only oxide). K + O₂ → KO₂ (superoxide).
Borax bead test: fused bead shows characteristic colour of metal oxide → qualitative analysis.
SiO₂ + HF → SiF₄ (HF attacks glass — why? F⁻ forms strong bond with Si). All other acids do NOT attack SiO₂.
Allotropes of carbon: diamond (sp³, hardest known), graphite (sp², electrical conductor), fullerene (C₆₀).
Markovnikov's rule: in addition to unsymmetrical alkene, H attaches to C with MORE hydrogens.
Electrophilic aromatic substitution: benzene ring attacks electrophile; ring is electron-rich due to delocalised π system.
Carcinogenicity: polycyclic aromatic hydrocarbons (PAHs) like benzene are carcinogenic — mention this in organic reactions.
Functional group order (acidity): COOH > phenol > H₂O > alcohol (higher the stability of conjugate base, stronger the acid).
Environmental chemistry: acid rain = SO₂ + H₂O → H₂SO₃; ozone depletion: Cl from CFC reacts with O₃ → O₂.