Chapter 3 Notes: Electronegativity & Bonding

Vocabulary-style flashcards covering electronegativity, bond types, VSEPR, hybridization, MO theory, and hydrogen bonding as presented in the notes.

Electronegativity

Power of an atom to attract shared electrons toward itself in a molecule.

Group Trend in Electronegativity

Electronegativity decreases down a group due to increasing atomic size and shielding, reducing attraction of bonding electrons by the nucleus.

Period Trend in Electronegativity

Electronegativity increases across a period due to increasing nuclear charge and decreasing atomic size, strengthening attraction for bonding electrons.

Ionic Bond (via EN difference)

An ionic bond is favored when the electronegativity difference between two atoms is large (roughly ΔEN > 1.8).

Covalent Bond (via EN difference)

A covalent bond is favored when the electronegativity difference is small (roughly ΔEN < 1.8).

Dipole Moment

A measure of bond polarity; μ = q × r; directed from the positive to the negative charge.

Polar Bond

A covalent bond with unequal sharing of electrons, producing a dipole moment.

Nonpolar Bond

A covalent bond where electrons are shared equally, resulting in negligible dipole moment.

Bond Energy

Energy required to break one mole of bonds in a substance under standard conditions; bond formation releases energy.

H–H Bond Energy

Bond energy of hydrogen gas (H–H) is about 436 kJ/mol.

Cl–Cl Bond Energy

Bond energy of chlorine gas (Cl–Cl) is about 243 kJ/mol.

Bond Length

Distance between two bonded atoms; typical values: C–C 1.54 Å, C=C 1.34 Å, C≡C 1.20 Å, C–H ~1.09 Å.

Bond Order

The number of chemical bonds between a pair of atoms; higher bond order yields shorter, stronger bonds.

VSEPR Theory

Valence Shell Electron Pair Repulsion; electron pairs around a central atom arrange to minimize repulsion, determining molecular geometry.

Electron Pair Geometry vs Molecular Geometry

Electron pair geometry considers all electron pairs; molecular geometry considers only the arrangement of bonded atoms.

Lone Pair

A nonbonding pair of valence electrons; occupies more space than a bonding pair and influences molecular shape.

Linear (AB2) Geometry

Two bonding electron pairs, zero lone pairs; bond angle 180°.

Bent (V‑shaped) Geometry

Molecule with two bonding pairs and two lone pairs (AB2E2) or similar; bond angle < 109.5° (e.g., H2O ~104.5°).

Trigonal Planar

Three bonding pairs, zero lone pairs (AB3); bond angles ~120°.

Tetrahedral

Four bonding pairs, zero lone pairs (AB4); bond angles ~109.5°.

Trigonal Pyramidal

Three bonding pairs and one lone pair (AB3E); bond angle ~107°.

Hybridization

Mixing of atomic orbitals to form new equivalent hybrid orbitals (sp, sp2, sp3) that explain molecular shapes.

sp3 Hybridization

Four equivalent sp3 hybrids; tetrahedral geometry; ~109.5° angles (e.g., CH4).

sp2 Hybridization

Three sp2 hybrids; trigonal planar geometry; ~120° angles (e.g., BF3).

sp Hybridization

Two sp hybrids; linear geometry; ~180° (e.g., C2H2).

Sigma Bond

Bond formed by end-to-end overlap along the internuclear axis; typically the first bond in a molecule.

Pi Bond

Bond formed by sideways overlap of p orbitals; exists in double and triple bonds in addition to a sigma bond.

Molecular Orbital Theory (MOT)

Theory that bonds arise from the combination of atomic orbitals to form molecular orbitals; electrons occupy these MOs.

Bonding Molecular Orbital (BMO)

Lower-energy MO with electron density between nuclei; contributes to a stronger bond.

Antibonding Molecular Orbital (ABMO)

Higher-energy MO with a node between nuclei; decreases bond strength.

Bond Order (MOT)

BO = (electrons in bonding MOs − electrons in antibonding MOs) / 2; positive values indicate a bond; higher BO means stronger bond.

Paramagnetism

Material with unpaired electrons; attracted by magnetic fields.

Diamagnetism

Material with all electrons paired; weakly repelled by magnetic fields.

Hydrogen Bonding

Intermolecular attraction where a hydrogen attached to N, O, or F forms a bond with a lone pair on another electronegative atom; raises boiling/melting points, affects solubility.

Solubility and Hydrogen Bonding

Alcohols dissolve in water because they can form hydrogen bonds; carboxylic acids dissolve due to –OH and C=O groups enabling hydrogen bonding.