Covalent bond

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A covalent bond is a chemical bond formed when atoms share electrons to make bonding pairs. By sharing, each atom can fill its outer shell and reach a stable electronic configuration. Covalent bonding is especially common in organic chemistry.

How covalent bonds work
- Bond strength and polarity depend on how similarly the atoms attract electrons (electronegativity). Similar electronegativities give nonpolar covalent bonds (like H–H); different electronegativities give polar covalent bonds (like H–Cl) when the geometry doesn’t cancel the dipoles.
- A single covalent bond uses one shared pair of electrons (a sigma, or σ, bond). A double bond adds a second shared pair (one sigma and one pi, or π, bond). A triple bond adds a third shared pair (one sigma and two pi bonds).
- Sigma bonds are made by head-on overlap of orbitals; pi bonds come from sideways overlap of p (or d) orbitals and are generally weaker than sigma bonds.

Types of covalent substances
- Molecules: individual units held together by covalent bonds (for example, H2, CH4, H2O).
- Molecular substances: weaker attractions between molecules give low boiling or melting points (like ethanol or iodine).
- Giant covalent networks: lots of atoms connected by covalent bonds in a lattice (diamond, graphite, quartz). These have high melting points and are usually hard and insulative.
- Polymers and biopolymers: long chains formed by repeating covalent bonds (polyethylene, proteins, carbohydrates).

Delocalization and resonance
- In some systems, electrons are shared over many atoms rather than confined to a single pair. This makes bonds more stable in compounds like benzene, a concept known as aromaticity. In benzene, six π electrons are spread over a ring, giving extra stability.

Unusual covalent bonding
- Three-center and four-electron bonds: some molecules share electrons among three or more atoms (for example, in diborane B2H6 there are three-center two-electron bonds). Expanded bonding can occur in some fluorides and other compounds (three-center four-electron bonds explain certain structures).
- Odd-electron bonds and radicals: molecules can have an unpaired electron, leading to reactive species and sometimes “half bonds.”
- Resonance and fractional bond orders: some molecules cannot be described by a single Lewis structure; several structures contribute, giving an average bond order (for example, in nitrate, NO3−).

Historical ideas and theories
- Early ideas focused on shared valence electrons. Lewis introduced electron dot structures to show how atoms share electrons to complete their outer shells.
- The term covalence was used in the early 20th century to describe the number of shared electron pairs.
- Valence Bond (VB) theory emphasizes localized bonds and resonance between structures; Molecular Orbital (MO) theory builds bonds from delocalized orbitals. VB is good for bond energies and mechanisms; MO helps explain spectra and ionization. Modern chemistry uses concepts from both theories.

In short, covalent bonding explains how atoms stick together by sharing electrons, forming a wide range of substances from simple molecules to giant networks, with a variety of bond types and degrees of electron sharing that scientists describe with different theoretical tools.