Subjects and lessons to explore.

Law

General Chemistry I · Chemical Bonding & Molecular Geometry

Valence Bond Theory and Hybridization

5 min read
SourcesReport an error
Want it in plain words first? Jump to Eli explains — the same idea, no jargon.
On this page 7 sections
  1. In 30 seconds
  2. Why this matters
  3. The college version
  4. Eli explains
  5. Key takeaway
  6. Study tools
  7. Sources & references

In 30 seconds

Valence bond theory pictures a covalent bond as the overlap of atomic orbitals, with a shared electron pair concentrated in the overlap region. To explain real molecular geometries, atomic orbitals are mixed (hybridized) into an equal number of hybrid orbitals. A sigma (σ) bond forms from head-on overlap of orbitals along the bond axis; pi (π) bonds form from side-by-side overlap of unhybridized p orbitals. The hybridization of an atom (sp, sp², sp³, sp³d, sp³d²) is set by the number of electron domains around it.

Why this matters

Hybridization explains the shapes and bond angles that plain atomic orbitals cannot. It accounts for why carbon forms four identical bonds in methane (sp³), why ethene is flat with 120° angles (sp²), and why acetylene is linear (sp). The σ/π distinction explains bond strength and rigidity: π bonds prevent rotation, which is why double bonds lock molecules into cis/trans shapes — a fact central to everything from fats to vision chemistry.

The college version

Key Ideas

  • Valence bond (VB) theory: a bond forms when atomic orbitals overlap and two electrons (one from each atom) occupy the overlap region.
  • Hybridization: mixing valence atomic orbitals into an equal number of equivalent hybrid orbitals pointing at the bonded atoms.
  • Sigma (σ) bond: head-on overlap along the internuclear axis; every single bond is a σ bond.
  • Pi (π) bond: side-by-side overlap of parallel p orbitals above and below the axis; weaker than σ.
  • Bond composition: single = 1 σ; double = 1 σ + 1 π; triple = 1 σ + 2 π.
  • Hybridization ↔ geometry: sp (2 domains, linear), sp² (3, trigonal planar), sp³ (4, tetrahedral), sp³d (5, trigonal bipyramidal), sp³d² (6, octahedral).

Equations and Variables

  • Number of hybrid orbitals = number of atomic orbitals mixed (sp = 2, sp² = 3, sp³ = 4, sp³d = 5, sp³d² = 6).
  • Rule of thumb: hybridization = number of electron domains (σ bonds + lone pairs) on the central atom. A double bond contributes one σ (its π part uses an unhybridized p orbital).

How It Works

  1. Count the electron domains on the central atom (σ bonds + lone pairs; each multiple bond contributes one σ).
  2. Mix that many atomic orbitals to make hybrid orbitals with the geometry that matches the domain count.
  3. Form σ bonds by head-on overlap of hybrid orbitals (or a 1s orbital on H) with other atoms.
  4. Any leftover unhybridized p orbitals overlap side-by-side to form π bonds in double and triple bonds.
  5. Lone pairs occupy hybrid orbitals, completing the geometry.

Worked Example

Determine the hybridization of the central atom in CH₄, BF₃, BeCl₂, and SF₆.

  1. CH₄: C has 4 domains (4 σ bonds) → sp³ (tetrahedral, 109.5°).
  2. BF₃: B has 3 domains (3 σ bonds) → sp² (trigonal planar, 120°).
  3. BeCl₂: Be has 2 domains (2 σ bonds) → sp (linear, 180°).
  4. SF₆: S has 6 domains (6 σ bonds) → sp³d² (octahedral, 90°).

Count σ and π bonds in ethene (C₂H₄, H₂C=CH₂): each carbon is sp² (3 domains: two C–H σ bonds + one C–C σ bond). The C=C double bond is 1 σ + 1 π; the π bond comes from side-by-side overlap of the leftover unhybridized p orbital on each carbon. Total: 5 σ bonds (4 C–H + 1 C–C) and 1 π bond.

Common Confusions

  • "A double bond is two σ bonds." — Wrong: a double bond is one σ plus one π; a triple is one σ plus two π.
  • "π bonds are stronger than σ bonds." — Wrong: π bonds (side-on overlap) are weaker than σ bonds.
  • "Hybridization is determined by the number of bonds only." — Wrong: it's the number of electron domains (σ bonds + lone pairs). NH₃ and H₂O are sp³ because of their lone pairs.
  • "Hybrid orbitals are real orbitals that existed before bonding." — They are a model (a mathematical mixing) used to describe the bonding; the atoms don't literally "pre-mix" orbitals.
Eli, the EliExplains learning guide

Eli explains

The same idea, in plain words

Explain it like I’m 10

Imagine two hands high-fiving. A σ bond is a straight-on high-five — solid, direct contact along the line between them. A π bond is two hands brushing sideways, overlapping only at the edges — weaker and more delicate. To make the right number of high-fives in the right directions, an atom "reshuffles" its available hands into hybrid positions first: two hands for a straight line (sp), three for a triangle (sp²), four for a pyramid (sp³). That reshuffling is hybridization. (The analogy captures overlap and directionality; real orbitals are quantum probability clouds, not hands, but the head-on vs. side-on overlap distinction is faithful.)

Key takeaways

  • VB theory: bonds = orbital overlap with a shared electron pair.
  • σ bond = head-on overlap; π bond = side-by-side p-orbital overlap.
  • Single bond = 1 σ; double = 1 σ + 1 π; triple = 1 σ + 2 π.
  • Hybridization by domain count: sp (2), sp² (3), sp³ (4), sp³d (5), sp³d² (6).
  • Geometry: sp linear, sp² trigonal planar, sp³ tetrahedral, sp³d trigonal bipyramidal, sp³d² octahedral.
  • Lone pairs occupy hybrid orbitals.
  • Examples: CH₄ sp³, BF₃ sp², BeCl₂/CO₂ sp, PCl₅ sp³d, SF₆ sp³d².
  • π bonds are weaker than σ bonds and block rotation.
  • Bond = orbital overlap (VB theory).
  • σ = head-on; π = side-by-side p overlap.
  • Single/double/triple = 1σ / 1σ+1π / 1σ+2π.
  • Hybridization = number of electron domains (σ bonds + lone pairs).
  • sp/sp²/sp³/sp³d/sp³d² → linear/trigonal planar/tetrahedral/trigonal bipyramidal/octahedral.
  • Lone pairs sit in hybrid orbitals.
  • π bonds block rotation → cis/trans isomerism.

Keep learning

Ready to build on this? Continue to the next lesson.

Practice General Chemistry I

This lesson has no separate scored set. Practice draws from the subject’s question bank.

Study tools & related lessonsYou’ll learn to · Related

You’ll learn to

  • Explain how valence bond theory describes covalent bonding as orbital overlap.
  • Distinguish sigma (σ) and pi (π) bonds.
  • Determine the hybridization of a central atom from its electron-domain count.
  • Relate hybridization to molecular geometry (sp, sp², sp³, sp³d, sp³d²).

Sources & references

  1. OpenStax, *Chemistry 2e*, "8.1 Valence Bond Theory."
  2. OpenStax, *Chemistry 2e*, "8.2 Hybrid Atomic Orbitals."
  3. OpenStax, *Chemistry 2e*, "8.3 Multiple Bonds."

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

Educational content only. It is not medical, legal or professional advice. Found an error? Tell us.