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Bonds are classified based on electronegativity difference and electron behavior. Ionic bonds (Δ EN > 1.7) involve electron transfer; covalent bonds (Δ EN < 1.7) involve sharing; metallic bonds involve delocalized electrons in a lattice. This classification predicts compound properties like melting point, solubility, and conductivity.
You have already studied how ionic bonds form through electron transfer and how covalent bonds form through electron sharing. The classification of bonds brings these two models together with a third — metallic bonding — and reveals that these are not three completely separate phenomena but rather points along a continuous spectrum determined by how atoms share or distribute their electrons.
The key variable is electronegativity difference (ΔEN) between the bonded atoms. When ΔEN is large (roughly above 1.7), one atom pulls electrons so strongly that they effectively transfer completely, creating oppositely charged ions held together by electrostatic attraction — an ionic bond. Sodium chloride is the classic example: sodium (EN ≈ 0.9) and chlorine (EN ≈ 3.2) differ by 2.3, so sodium gives up its valence electron entirely. When ΔEN is small (below about 1.7), neither atom dominates, and electrons are shared between the nuclei — a covalent bond. The sharing may be equal (as in H₂ or Cl₂, where ΔEN = 0) or unequal (as in H–Cl, where chlorine pulls the shared pair closer, creating a polar covalent bond). The 1.7 threshold is a guideline, not a sharp boundary — bonding character transitions gradually from purely covalent to purely ionic.
Metallic bonding represents a third arrangement that appears when atoms of low electronegativity pack together. Instead of transferring electrons to a partner or sharing them in localized pairs, metal atoms release their valence electrons into a communal "sea" that pervades the entire lattice. Each metal cation sits in a regular array, surrounded by freely mobile electrons that belong to no single atom. This delocalized electron model explains why metals conduct electricity (electrons move freely), are malleable (layers of cations can slide without breaking bonds), and have luster (free electrons absorb and re-emit light across the visible spectrum).
The real power of bond classification is its predictive reach. Once you identify the bond type, you can anticipate macroscopic properties without memorizing them individually. Ionic compounds form crystalline lattices with high melting points because every ion is locked in place by strong electrostatic forces in all directions; they dissolve in polar solvents and conduct electricity when melted or dissolved because the ions become free to move. Covalent compounds form discrete molecules with lower melting points because the forces between molecules (intermolecular forces) are much weaker than the bonds within them; they are often poor conductors because they have no free charges. Metals conduct in the solid state, are ductile, and have moderate to high melting points depending on how many electrons each atom contributes to the sea. Recognizing that these property patterns flow directly from electron behavior — transferred, shared, or delocalized — is the central insight of bond classification.