AS and A Level Chemistry Revision Notes

Subatomic Particles: Atoms comprise protons ( $+1$ charge, $1$ relative mass), neutrons ( $0$ charge, $1$ relative mass), and electrons ( $-1$ charge, negligible mass).
Relative Masses:
- Relative atomic mass ( $A_r$ ): Average mass of an atom relative to $\frac{1}{12}$ of the mass of a carbon-12 atom.
- Relative molecular mass ( $M_r$ ): Total mass of a molecule relative to the same carbon-12 standard.
The Mole: Defined by the Avogadro constant ( $6.02 \times 10^{23}$ ). It represents the number of atoms in the relative atomic mass of a substance expressed in grams.
Stoichiometry Calculations:
- Empirical Formula: The simplest whole-number ratio of atoms in a compound.
- Molecular Formula: The actual number of atoms of each element in a molecule.
- Gas Volume: One mole of any gas occupies $24\,dm^3$ at room temperature ( $25^\circ\text{C}$ ) and $1\,atm$ pressure.
- Concentration: Molar concentration is measured in $mol\,dm^{-3}$ .

Orbitals and Shells: Electrons occupy energy levels divided into sub-shells ( $s, p, d$ ). The filling sequence follows the order: $1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p$ .
Ionisation Energy (IE): The energy needed to remove one mole of electrons from one mole of gaseous atoms to form gaseous cations.
- Trends: General increase across a period due to increasing nuclear charge and constant shielding; decrease down a group as the distance from the nucleus increases.
- Successive IEs: Large jumps in IE indicate the removal of an electron from a new inner shell.

Bond Types:
- Ionic: Electrostatic attraction between oppositely charged ions in a giant lattice.
- Covalent: Shared pairs of electrons; can be coordinate (dative) if both electrons originate from one atom.
- Metallic: Positive metal ions in a "sea" of delocalised electrons.
VSEPR Theory: Molecular shapes (Linear, Trigonal Planar, Tetrahedral, etc.) are determined by the repulsion between bonding and lone pairs of electrons.
Electronegativity: The power of an atom to attract a bonding pair of electrons; increases across a period and decreases down a group.
Intermolecular Forces:
- Van der Waals Forces: Weakest attractions caused by instantaneous dipoles.
- Permanent Dipole-Dipole: Attractions between polar molecules.
- Hydrogen Bonding: A strong dipole-dipole interaction involving hydrogen bonded to $N, O,$ or $F$ .

Ideal Gases: Defined by the equation $pV = nRT$ . Real gases deviate at high pressures and low temperatures where molecular volume and attractions become significant.
Enthalpy Changes (\Delta H):
- Standard Enthalpy of Formation (\Delta H_f^\theta): Formation of $1\,mol$ of a compound from its elements under standard conditions ( $298\,K, 101\,kPa$ ).
- Standard Enthalpy of Combustion (\Delta H_c^\theta): Complete combustion of $1\,mol$ of a substance in excess oxygen.
Hess’s Law: The total enthalpy change for a reaction is independent of the route taken.
Born–Haber Cycles: Used to determine lattice energy by combining various enthalpy changes (atomisation, ionisation, electron affinity).

Redox Processes: Oxidation involves loss of electrons; reduction involves gain.
Standard Electrode Potentials (E^\theta): Measured using a standard hydrogen electrode ( $E^\theta = 0.00\,V$ ). Positive values indicate better oxidising agents.
Faraday’s Constant: The charge per mole of electrons ( $F \approx 96500\,C\,mol^{-1}$ ).
Dynamic Equilibrium: Occurs when the rates of forward and reverse reactions are equal.
- Le Chatelier’s Principle: Systems shift to counteract changes in concentration, pressure, or temperature.
- Equilibrium Constants: $K_c$ (concentration) and $K_p$ (partial pressure) vary with temperature.
Acids and Bases: Br nsted–Lowry theory defines acids as proton donors and bases as proton acceptors.
- $pH = -\log[H^+]$ ; $K_w = [H^+][OH^-] = 10^{-14}$ at $25^\circ\text{C}$ .
- Buffer Solutions: Resist pH changes upon addition of small amounts of acid or alkali.

Collision Theory: Reactions require collisions with energy $\ge$ Activation Energy ( $E_a$ ).
Boltzmann Distribution: Describes the range of molecular energies; higher temperatures increase the proportion of molecules with $E \ge E_a$ .
Reaction Orders: Determined experimentally ( $0, 1, 2$ order). Half-life ( $t_{1/2}$ ) is constant for first-order reactions.
Period 3 Periodicity: Atomic radius decreases across the period; melting points peak at Silicon (giant covalent) before dropping for Phosphorus, Sulfur, and Chlorine (simple molecular).

Functional Groups: Alkanes (saturated), Alkenes (unsaturated, $C=C$ ), Halogenoalkanes ( $C-X$ ), Alcohols ( $OH$ ), Aldehydes ( $CHO$ ), Ketones ( $CO$ ), Carboxylic Acids ( $COOH$ ).
Reaction Mechanisms:
- Free-radical substitution: Initiation, propagation, and termination steps.
- Electrophilic Addition: Typical for alkenes (e.g., bromination).
- Nucleophilic Substitution: $S_N 1$ (two steps via carbocation) or $S_N 2$ (one transition state step).
Isomerism:
- Structural: Chain, position, and functional group isomers.
- Stereoisomerism: cis-trans ( $geometric$ ) and optical (involves a $chiral$ centre with four different groups).

Proteins: Polymers of 2-amino acids joined by peptide bonds. Structure is defined as primary (sequence), secondary ( $\alpha$ -helix, $\beta$ -sheet), and tertiary (folding).
DNA: Double helix formed by nucleotides (sugar, phosphate, and bases: $A, T, C, G$ ). Base pairing is complementary ( $A-T, C-G$ ).
Analytical Techniques:
- Electrophoresis: Separation based on charge and size.
- NMR Spectroscopy: Identifies distinct chemical environments and spin-spin splitting ( $n+1$ rule).
- Mass Spectrometry: Determines relative atomic masses and identifies molecular fragments.
- Chromatography: Separation based on partition or adsorption between stationary and mobile phases.

Ionisation Energy Variations: Discussion regarding why specific elements like Boron or Oxygen show lower than expected energy due to sub-shell transitions and electron-pair repulsion.
Enthalpy Calculations: Examiners emphasize the standard definition of formation, including the requirement for reagents to be in their standard states.
Organic Structure Identification: Strategies for using NMR and mass spectrometry data to deduce molecular structures, specifically looking at chemical shifts for groups like $-CH_3, -OH,$ and aromatic rings.
Solubility and Lattice Energy: Analysis of how cation size and charge density affect the stability of Group II carbonates and sulfates.

Would you like the summary of the next section, such as the detailed qualitative analysis tests or more A2 organic mechanisms?