Chemical Bonding & Molecular Structure

Vocabulary flashcards covering key terms and definitions from the lecture on chemical bonding, molecular structure, hybridization, intermolecular forces, and molecular orbital theory.

Chemical bond

A very strong force of attraction that binds two atoms or ions together.

Ionic (Electrovalent) bond

Bond formed by complete transfer of electrons from an electropositive atom to an electronegative atom.

Factors favoring ionic bond

Low ionization energy of metal, high electron affinity of non-metal, and high lattice energy of the resulting crystal.

Characteristics of ionic compounds

Ionic lattice, usually solid, soluble in polar solvents, conduct electricity in solution, high melting/boiling points, non-directional bonds.

Covalent bond

Bond produced by equal sharing of electron pairs between two atoms.

Single covalent bond

One shared pair of electrons between two atoms (e.g., H–H).

Double covalent bond

Two shared electron pairs between atoms (e.g., O=C=O).

Triple covalent bond

Three shared electron pairs between atoms (e.g., N≡N).

Coordinate (dative) bond

Covalent bond in which the shared pair is donated by one atom (donor) and accepted by the other (acceptor).

Odd-electron bond

Bond present in molecules/ions containing an unpaired electron, e.g., NO, NO₂, ClO₂.

Metallic bond

Attraction between positive metal ions and a ‘sea’ of delocalized electrons in a metal lattice.

Bond length

Average distance between nuclei of two bonded atoms; increases with atomic size, decreases with multiple bonding and higher s-character.

Bond energy (bond strength)

Energy released on bond formation or required for bond breaking; increases from single to triple bonds and with fewer lone pairs.

Polarization of ionic bond

Distortion of anion electron cloud by a cation, reducing ionic character.

Fajan’s rule

Smaller, highly charged cation and larger, highly charged anion cause greatest polarization and covalent character.

Polarization of covalent bond

Shift of shared electrons toward the more electronegative atom, giving partial ionic character.

Dipole moment (μ)

Product of charge (q) and bond length (d); vector quantity expressed in Debye; indicates bond polarity.

Formal charge

FC = valence electrons – lone-pair electrons – ½(shared electrons); helps choose the best Lewis structure.

Hybridization

Mixing of atomic orbitals of similar energy on the same atom to form equivalent hybrid orbitals.

Rules of hybridization

Same-atom orbitals mix; number of hybrids equals number of orbitals mixed; hybrids form σ-bonds; arranged for maximum separation.

sp hybridization

Linear geometry, 180° bond angle; e.g., BeF₂, CO₂, C₂H₂.

sp² hybridization

Trigonal planar geometry, 120° bond angle; e.g., BF₃, C₂H₄, SO₃.

sp³ hybridization

Tetrahedral electron geometry (109°28′); 0–2 lone pairs give shapes: tetrahedral, pyramidal (NH₃), V-shaped (H₂O).

dsp² hybridization

Square planar geometry, 90° angles; e.g., [Ni(CN)₄]²⁻, [PtCl₄]²⁻.

sp³d hybridization

Trigonal bipyramidal electron geometry; shapes vary with lone pairs: TBP (PCl₅), see-saw (SF₄), T-shape (ClF₃), linear (XeF₂).

sp³d² hybridization

Octahedral electron geometry; shapes: octahedral (SF₆), square pyramidal (ICl₅), square planar (XeF₄).

Sigma (σ) bond

Bond formed by end-to-end overlap with electron density along the internuclear axis; allows free rotation, defines molecular shape.

Pi (π) bond

Bond formed by sideways overlap of p-orbitals above/below the internuclear axis; prohibits rotation and accompanies a σ-bond.

Dipole–dipole attraction

Intermolecular force between opposite poles of two permanent dipoles.

Hydrogen bond

Strong dipole–dipole attraction between H attached to F, O, or N and a lone pair on F, O, or N of another molecule.

London (dispersion) forces

Attractions between instantaneous or induced dipoles; present in all molecules, dominant in non-polar species.

Ion-dipole attraction

Force between an ion and the oppositely charged pole of a polar molecule.

Molecular Orbital Theory (MOT)

Approach treating electrons as delocalized over the whole molecule, forming molecular orbitals that extend over all nuclei.

Bonding molecular orbital

MO formed by constructive interference (in-phase overlap) of atomic orbitals; lower energy, increases electron density between nuclei.

Antibonding molecular orbital

MO formed by destructive interference (out-of-phase overlap); higher energy, contains a node between nuclei, labeled with an asterisk (*).

Bond order (MOT)

(Number of bonding electrons – number of antibonding electrons)/2; indicates bond strength and length.

Stability criterion in MOT

A molecule is stable if bonding electrons (Nb) exceed antibonding electrons (Na); zero or negative bond order predicts non-existence.