AP Chemistry Unit 2 Essentials: Key Terms & Big Ideas

Big Idea 1

Electronegativity difference predicts bond type

Whether two atoms form an ionic, covalent, or metallic bond comes down to how unevenly they share electrons. Large electronegativity differences (metal + nonmetal) cause full electron transfer — ionic bonding. Small or zero differences between nonmetals cause electron sharing — covalent bonding. Metal atoms pooling their valence electrons into a shared "sea" produces metallic bonding. Once you can estimate electronegativity difference, you can predict bond type before doing anything else.

Electronegativity Bond Type Coulomb's Law

Big Idea 2

Lewis structures are a model, and formal charge tells you which model is best

A Lewis structure is chemistry's shorthand for showing where electrons are most likely to be found. When more than one valid structure can be drawn for a molecule (resonance), formal charge — a bookkeeping tool, not a real charge — helps you decide which structure (or combination of structures) is the most realistic representation of the true, delocalized electron distribution.

Lewis Structures Resonance Formal Charge

Big Idea 3

Electron domains, not atoms, determine molecular shape

VSEPR theory works because electron domains — bonding pairs and lone pairs alike — repel each other and arrange themselves as far apart as possible. Lone pairs take up more space than bonding pairs, which is why water bends more than its electron-domain geometry alone would suggest. Hybridization (sp, sp², sp³) is simply the orbital-level explanation for why these domain geometries occur, and it sets up which bonds are sigma (head-on overlap) versus pi (side-by-side overlap) — the language Unit 3 builds directly on.

VSEPR Hybridization Molecular Geometry

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Ionic bond

A bond formed by the electrostatic attraction between oppositely charged ions, typically created when a metal transfers electrons to a nonmetal.

Bonding Types

Covalent bond

A bond formed when two atoms (usually nonmetals) share one or more pairs of electrons.

Bonding Types

Metallic bond

A bond in which metal atoms share a "sea" of delocalized valence electrons that move freely throughout the structure, explaining conductivity and malleability.

Bonding Types

Bond length

The average distance between the nuclei of two bonded atoms. Shorter bonds are generally stronger (e.g., triple bonds are shorter than double or single bonds between the same atoms).

Bond Energy

Bond energy

The energy required to break one mole of a particular covalent bond in the gas phase. Higher bond energy means a stronger, harder-to-break bond.

Bond Energy

Potential energy diagram

A graph showing potential energy versus distance between two atoms; the minimum point corresponds to the most stable bond length.

Bond Energy

Lattice energy

The energy released when gaseous ions combine to form one mole of an ionic solid. Increases with higher ionic charge and decreases with larger ionic radius.

Ionic Solids

Coordination number

The number of ions immediately surrounding a given ion in a crystal lattice.

Ionic Solids

Electron sea model

A model of metallic bonding describing valence electrons as delocalized and free to move throughout a lattice of metal cations, explaining conductivity, malleability, and ductility.

Metals & Alloys

Alloy

A mixture of a metal with one or more other elements (often other metals), which changes properties like hardness, strength, or corrosion resistance.

Metals & Alloys

Malleability / ductility

Malleability is the ability to be hammered into thin sheets; ductility is the ability to be drawn into wires. Both result from delocalized electrons allowing metal atoms to slide past one another without breaking bonds.

Metals & Alloys

Lewis structure

A diagram showing the bonding pairs and lone pairs of valence electrons in a molecule or ion using dots and lines.

Lewis Structures

Octet rule

The tendency of main-group atoms to gain, lose, or share electrons to achieve eight valence electrons (a stable noble-gas configuration).

Lewis Structures

Lone pair

A pair of valence electrons on an atom that is not shared in a bond. Lone pairs occupy more space than bonding pairs and influence molecular shape.

Lewis Structures

Formal charge

A bookkeeping charge assigned to an atom in a Lewis structure, calculated as (valence electrons) − (nonbonding electrons) − (bonding electrons ÷ 2). Used to evaluate which of several possible Lewis structures is most reasonable.

Resonance

Resonance structure

One of two or more valid Lewis structures for the same molecule that differ only in the position of electrons (not atoms). The real structure is a blend (resonance hybrid) of all contributing structures.

Resonance

VSEPR theory

Valence Shell Electron Pair Repulsion theory — predicts molecular shape based on the principle that electron domains (bonding and nonbonding) arrange themselves to minimize repulsion.

VSEPR

Electron-domain geometry vs. molecular geometry

Electron-domain geometry counts all electron domains (bonding + lone pairs); molecular geometry describes only the positions of bonded atoms. They differ whenever lone pairs are present (e.g., water is tetrahedral electron-domain but bent molecular geometry).

VSEPR

Hybridization

The mixing of atomic orbitals (s, p, sometimes d) into new hybrid orbitals (sp, sp², sp³) that match the geometry predicted by VSEPR and provide better orbital overlap for bonding.

Hybridization

Sigma (σ) bond

A covalent bond formed by the head-on overlap of orbitals along the axis connecting two nuclei. Every single, double, and triple bond contains exactly one sigma bond.

Hybridization

Pi (π) bond

A covalent bond formed by the side-by-side overlap of unhybridized p orbitals above and below the bond axis. Double bonds have one pi bond; triple bonds have two.

Hybridization