Sigma and Pi Bonds & Hybrid Orbitals – Vocabulary

Vocabulary flashcards summarizing the key bonding and hybridization terms discussed in Unit 5 – Part 6 lecture.

Covalent Bond

A bond formed when two atoms share a pair of electrons—one electron contributed by each atom.

Sigma (σ) Bond

The stronger type of covalent bond produced by direct, head-on overlap of orbitals along the line connecting two nuclei.

Pi (π) Bond

A weaker covalent bond created by lateral (sideways) overlap of parallel p orbitals; electron density is above and below the internuclear axis.

Valence Bond Theory

Bonding model in which covalent bonds arise from overlap of atomic or hybrid orbitals containing unpaired electrons; basis for VSEPR shapes.

Molecular Orbital Theory

Alternative, more complex bonding model that describes electrons in molecular orbitals spread over the whole molecule and explains properties such as magnetism.

Hybridization

The process of mixing atomic orbitals on one atom to form new, equivalent hybrid orbitals oriented for optimal bonding.

Hybrid Orbital

An orbital produced during hybridization; equal in energy and shape to others in its set and used for σ bonds or lone pairs.

sp3 Hybrid Orbital

Hybrid made from one s and three p orbitals; forms four equivalent orbitals arranged 109.5° apart (tetrahedral).

sp2 Hybrid Orbital

Hybrid made from one s and two p orbitals; forms three equivalent orbitals arranged 120° apart (trigonal planar).

sp Hybrid Orbital

Hybrid made from one s and one p orbital; forms two equivalent orbitals arranged 180° apart (linear).

Electron Domain

A region around an atom where electrons (bonding or lone pairs) are likely found; determines molecular geometry and required hybrid orbitals.

Tetrahedral Geometry

Shape with four electron domains around a central atom at 109.5° angles; associated with sp3 hybridization.

Trigonal Planar Geometry

Shape with three electron domains around a central atom at 120° angles; associated with sp2 hybridization.

Linear Geometry

Shape with two electron domains around a central atom at 180°; associated with sp hybridization.

Single Bond

A bond consisting of one σ bond only.

Double Bond

A bond consisting of one σ bond plus one π bond.

Triple Bond

A bond consisting of one σ bond plus two π bonds.

Relative Strength (σ vs π)

σ bonds are stronger than π bonds because head-on overlap places electron density directly between nuclei, maximizing attraction.

Head-on Overlap

Orbital interaction along the internuclear axis that forms a σ bond.

Lateral (Sideways) Overlap

Parallel p-orbital interaction above and below the internuclear axis that forms a π bond.

p Orbital in π Bond

Unhybridized p orbital on each bonded atom that overlaps sideways to create a π bond.

s Orbital Contribution

Hydrogen’s 1s orbital overlaps with a hybrid orbital on another atom to make a σ bond.

Lone (Nonbonding) Pair

A pair of valence electrons localized on one atom; occupies a hybrid orbital and influences geometry.

Promotion (s→p)

Energy-requiring step where an s electron is excited to an empty p orbital so the atom attains enough unpaired electrons for bonding before hybridization.