A bond is an energy story
Imagine two hydrogen atoms approaching from far apart. Each nucleus attracts electrons, while nuclei repel nuclei and electrons repel electrons. At the right separation, the attractions win enough to make the joined arrangement lower in energy than the separated atoms. Move the nuclei too close and repulsion rises sharply.
That energy minimum is why a bond has a preferred length. Breaking a bond requires energy; energy is released only when the new bonds formed are more stabilising than the bonds that were broken.
Different electron distributions, not three sealed worlds
Covalent, ionic and metallic bonding are limiting pictures. Real bonding can sit between them, especially when electron density is unevenly shared.
Covalent
Electron density is concentrated between nuclei. In H₂, sharing stabilises the pair relative to two separated H atoms.
Ionic solid
Oppositely charged ions attract throughout an extended lattice. Solid NaCl is not a pile of isolated NaCl molecules.
Metallic
Valence electrons are delocalised across many atomic centres, helping explain why copper conducts while remaining a solid.
Electronegativity tilts a bond; geometry decides the molecule
In H–Cl, chlorine pulls the bonding electron density toward itself, so the bond has partial charges: Hδ+–Clδ−. That shift is gradual. There is no universal electronegativity-difference switch that suddenly turns a bond from “covalent” into “ionic”.
And a polar bond does not automatically make a polar molecule. Carbon dioxide has two polar C=O bonds, but its linear geometry makes their dipoles cancel.
The model earns its keep when it predicts behaviour
Does not conduct as a solid; molten salt conducts because ions can move.
Discrete H₂O molecules interact strongly through hydrogen bonding.
Conducts electricity in the solid state.
Very hard because breaking the structure means disrupting strong bonds throughout the network.
Lewis structures are maps, not photographs
Lewis structures are excellent for counting valence electrons, seeing formal charges and proposing bonding patterns. They deliberately leave out much of the quantum picture. Resonance makes that limitation obvious: benzene is not flickering between two drawings; the real electron distribution is represented by the resonance hybrid.