Classes of chemical bonds and intermolecular forces: ionic, covalent, metallic, hydrogen bonding, van der Waals; geometry rules (VSEPR) and molecular-orbital theory. Emerges from quantum mechanics (overlap integrals, Pauli exclusion) and…
chemical-bonding
Ionic bond
Electrostatic attraction between oppositely charged ions formed by full electron transfer, typically between a metal and a non-metal with…
Covalent bond
Bond formed by shared electron pairs between atoms with similar electronegativity. Quantum-mechanical in origin (wavefunction overlap,…
Metallic bond
Non-localised bonding in metals: valence electrons form a delocalised 'sea' shared among positively-charged ion cores, explaining…
Hydrogen bond
Partially-electrostatic, partially-covalent interaction between an H atom covalently bound to a highly electronegative atom (N, O, F) and a…
Van der Waals forces
Weak, distance-dependent attractive forces between neutral atoms and molecules: London dispersion (induced-dipole/induced-dipole), Debye…
VSEPR theory
Valence-shell electron-pair repulsion theory: the geometry of a small molecule is determined by minimising repulsion between electron pairs…
Molecular-orbital theory
Quantum-mechanical description of chemical bonding: linear combinations of atomic orbitals form delocalised molecular orbitals occupied…
Resonance
Description of delocalisation by drawing multiple Lewis structures connected by arrows. No individual resonance structure exists…
Orbital hybridization (sp, sp², sp³)
Linear combination of atomic s- and p-orbitals on a central atom to match observed molecular geometry: sp (linear), sp² (trigonal planar),…
Dipole moment (μ)
Vector sum of bond dipole moments in a molecule. Nonzero μ implies IR-active vibrations; symmetry-imposed cancellation (CO₂, CCl₄) gives μ…
Lone pair
Non-bonding valence-electron pair on an atom. Centrally important in VSEPR (lone pairs take more angular space than bonds), in Lewis…
Lewis structure (electron-dot diagram)
A 2D representation of a molecule showing valence electrons as dots and bonds as lines, with bonding pairs shared between atoms and…
Octet rule
Main-group atoms tend to gain, lose, or share electrons so as to complete a valence shell of 8 electrons (matching the nearest noble gas).…
Formal charge
A bookkeeping charge assigned to each atom in a Lewis structure, computed as FC = V − L − ½B where V = valence electrons, L = lone-pair…
Sigma (σ) bond
A covalent bond with cylindrical symmetry about the internuclear axis, formed by head-on (axial) overlap of atomic orbitals — s–s, s–p, or…
Pi (π) bond
A covalent bond formed by sideways (lateral) overlap of parallel p-orbitals above and below the internuclear axis. Weaker than a σ-bond and…
Valence bond theory & hybridization
Pauling's sp, sp², sp³ hybrids; resonance structures; Heitler–London H₂ treatment; complements MO theory.
Walsh diagrams (AH₂, AH₃)
MO energies vs molecular geometry; predicts bent/linear AH₂ from electron count; BeH₂ linear, H₂O bent.
Bent (banana) bonds
Alternative to σ+π: cyclopropane's strained C-C-C; two banana bonds; Pauling equivalence to classical picture.
3c-2e bonding (boranes, bridging H)
B₂H₆ bridging H via 3 atoms sharing 2 electrons; diborane, carboranes; Wade's rules for cluster counts.
Hyperconjugation
σ(C-H) donation into adjacent π*/empty p-orbital; stabilizes carbocations, alkenes; anomeric effect in sugars.
Hückel aromaticity (4n+2)
Planar, cyclic, conjugated with 4n+2 π electrons → aromatic stability; antiaromatic 4n; NICS probe.
Antiaromaticity (4n)
4n π electron cyclic systems destabilized; cyclobutadiene distorts to rectangular to avoid degeneracy; pentalene, biphenylene.
Möbius aromaticity (4n)
Topologically-twisted conjugated π systems aromatic at 4n electrons (inverted Hückel); [12]annulene experiments.
Atoms-in-molecules (Bader QTAIM)
Topology of ρ(r); bond critical points, atomic basins via zero-flux surfaces; rigorous definition of atomic properties.
Electron localization function (ELF)
η(r) ∈ [0,1] identifies regions of paired electrons; visualizes lone pairs, bonds, cores; complements QTAIM.
Natural bond orbitals (NBO)
Localized orbitals from 1-RDM diagonalization; bonding/anti-bonding transferable; donor-acceptor hyperconjugation analysis.
Charge-transfer complexes
Donor-acceptor π-complexes with intense CT absorption; Mulliken theory; I₂-benzene, tetracyanoethylene.
Halogen, chalcogen, pnictogen bonds
σ-hole directional interactions R-X···Y beyond H-bonding; crystal engineering and drug design applications.
Metallic bonding (free electron model)
Positive cores in electron 'sea'; Sommerfeld/Drude; tight-binding bands; explains luster, conductivity, malleability.
Hypervalent 3c-4e bonding
SF₆, PF₅, XeF₂: described without d-orbitals via 3-center-4-electron; MO treatment of hypervalent molecules.
Hoffmann isolobal analogy
Fragments with similar frontier orbitals (e.g., CH₃ ≈ Mn(CO)₅) show analogous bonding; bridges organic/organometallic.
Lewis structures & extended octet
Octet rule plus exceptions (BF₃ sextet, SF₆ 12e⁻); formal charges; resonance averaging; VSEPR shape prediction.