Last reviewed 16 Sept 2026 · 8 min read
Fundamental particles
| Particle | Charge | Relative mass | Location |
|---|---|---|---|
| Proton | +1 ( C) | ≈ 1 u | Nucleus |
| Neutron | 0 | ≈ 1 u | Nucleus |
| Electron | −1 | ≈ 1/1836 u | Around nucleus |
- Atomic number Z = number of protons; mass number A = protons + neutrons.
- Isotopes — same Z, different A (¹²C, ¹⁴C); isobars — same A, different Z.
Atomic models (summary)
Dalton (indivisible atoms) → Thomson (plum pudding) → Rutherford (nuclear atom, alpha scattering) → Bohr (quantised orbits, hydrogen spectrum) → quantum mechanical model (Schrödinger equation — electrons in orbitals, regions of high probability).
Quantum numbers
| Quantum number | Symbol | Values | Describes |
|---|---|---|---|
| Principal | n | 1, 2, 3, … | Shell, size and energy |
| Azimuthal (angular momentum) | l | 0 to n − 1 (s = 0, p = 1, d = 2, f = 3) | Subshell, shape of orbital |
| Magnetic | −l to +l | Orientation of orbital (number of orbitals = 2l + 1) | |
| Spin | +½ or −½ | Electron spin |
- Maximum electrons: in a subshell → s = 2, p = 6, d = 10, f = 14; in a shell .
- Orbital shapes: s — spherical; p — dumbbell; d — cloverleaf (mostly).
Rules for electron filling
- Aufbau principle — electrons occupy orbitals in order of increasing energy ((n + l) rule): 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p …
- Pauli exclusion principle — no two electrons in an atom have the same set of four quantum numbers (an orbital holds at most two electrons with opposite spins).
- Hund's rule of maximum multiplicity — electrons occupy degenerate orbitals singly with parallel spins before pairing.
Examples: Na (11): 1s² 2s² 2p⁶ 3s¹; Ca (20): [Ar] 4s²; Fe (26): [Ar] 3d⁶ 4s²; Cr (24): [Ar] 3d⁵ 4s¹ and Cu (29): [Ar] 3d¹⁰ 4s¹ (exceptions — extra stability of half-filled and fully filled d subshells).
Periodic table and trends
- Modern periodic law: properties are periodic functions of atomic number (Moseley).
- Groups (vertical, 18) — similar valence configurations; periods (horizontal, 7).
- Blocks: s (groups 1–2), p (13–18), d (transition metals), f (lanthanides, actinides).
| Property | Across a period (left → right) | Down a group |
|---|---|---|
| Atomic radius | Decreases (increasing nuclear charge) | Increases (new shells) |
| Ionisation energy | Generally increases | Decreases |
| Electron affinity | Generally becomes more negative (more energy released) | Generally less negative |
| Electronegativity | Increases (fluorine highest, ≈ 4.0 Pauling) | Decreases |
| Metallic character | Decreases | Increases |
- Noble gases (group 18) — stable filled shells, very low reactivity.
- Cations are smaller than parent atoms; anions larger.
Chemical bonding
Atoms bond to achieve stable configurations (often an octet) and lower energy.
| Bond | Formation | Properties | Examples |
|---|---|---|---|
| Ionic (electrovalent) | Transfer of electrons between metal and non-metal (large electronegativity difference) | Crystalline solids, high melting points, conduct when molten/dissolved, soluble in polar solvents | NaCl, CaO, MgO, CaCl₂ |
| Covalent | Sharing of electron pairs between non-metals | Molecules; lower melting points (except network solids like diamond, quartz), poor conductors | H₂, H₂O, CH₄, CO₂, SiO₂ (network) |
| Coordinate (dative) | Shared pair donated by one atom | Complexes | NH₄⁺, H₃O⁺, metal complexes |
| Metallic | Positive ions in a "sea of delocalised electrons" | Conductivity, malleability, ductility, lustre | Fe, Cu, Al |
| Hydrogen bond | Attraction between H attached to F, O or N and another electronegative atom | Relatively strong intermolecular force → high boiling point of water, ice structure (less dense), properties of clays and cement paste water | H₂O, NH₃, HF |
| van der Waals forces | Weak dipole–dipole and dispersion forces | Weakest; noble gas liquefaction, adsorption | Between non-polar molecules |
Relative strength: covalent/ionic > metallic (varies) > hydrogen bond > van der Waals.
Polarity and dipole moment
- A covalent bond between atoms of different electronegativity is polar (partial charges).
- Dipole moment (debye); a molecule may be non-polar despite polar bonds if symmetric (CO₂, CH₄, CCl₄), while bent H₂O and pyramidal NH₃ are polar.
Molecular shapes — VSEPR theory and hybridisation
VSEPR: electron pairs around a central atom repel and arrange to minimise repulsion; lone pair–lone pair > lone pair–bond pair > bond pair–bond pair repulsion.
| Electron pairs (bonding + lone) | Hybridisation | Geometry | Bond angle | Examples |
|---|---|---|---|---|
| 2 + 0 | sp | Linear | 180° | CO₂, BeCl₂, C₂H₂ |
| 3 + 0 | sp² | Trigonal planar | 120° | BF₃, C₂H₄, CO₃²⁻ |
| 4 + 0 | sp³ | Tetrahedral | 109.5° | CH₄, SiO₄⁴⁻ (silicates) |
| 3 + 1 | sp³ | Trigonal pyramidal | ≈ 107° | NH₃ |
| 2 + 2 | sp³ | Bent (angular) | ≈ 104.5° | H₂O |
| 5 + 0 | sp³d | Trigonal bipyramidal | 90°, 120° | PCl₅ |
| 6 + 0 | sp³d² | Octahedral | 90° | SF₆; Al–O octahedra in clays |
- Sigma (σ) bonds — head-on overlap (single bonds); pi (π) bonds — sideways overlap (in double and triple bonds).