Table salt dissolves instantly in water and conducts electricity once it does. Diamond is one of the hardest materials known and doesn't conduct electricity at all. Copper wire is soft, malleable, and conducts electricity brilliantly even as a solid. All three are held together by atoms bonded to each other — and the reason they behave so differently comes down to which of chemistry's three basic bonding strategies each one uses, and specifically what happens to the electrons involved.
Ionic bonding: a complete transfer, then attraction
Ionic bonds form when one atom essentially gives up one or more electrons entirely to another atom, rather than sharing them. This typically happens between a metal, which readily loses electrons, and a non-metal, which readily gains them — sodium giving up an electron to chlorine, for instance, to form table salt. The result is two oppositely charged ions, held together purely by electrostatic attraction, arranged into a rigid, repeating crystal lattice. That lattice structure is why ionic compounds tend to be brittle solids with high melting points, and why they conduct electricity only once the ions are freed to move — dissolved in water, or melted — rather than as a solid, where the ions are locked in place.
Covalent sharing and the metallic electron sea
Covalent bonds, by contrast, involve atoms sharing electron pairs between them rather than one atom fully surrendering electrons to another — this is the bonding pattern that holds carbon atoms together in diamond, each carbon covalently bonded to four neighbours in an extraordinarily rigid three-dimensional lattice, which is exactly why diamond is so hard and, with no free-moving charged particles available, doesn't conduct electricity. Metallic bonding takes yet a third approach: metal atoms release their outer electrons into a shared, delocalised "sea" that isn't attached to any single atom or pair of atoms, with the resulting positive metal ions held together by their collective attraction to that surrounding electron sea. Because those electrons are free to move throughout the entire structure, metals conduct electricity and heat exceptionally well, and because the metal ions can shift position relative to each other without breaking the bonding (the electron sea simply reorganises around them), metals are malleable and ductile in a way ionic and covalent solids generally aren't.
What we're still unsure about
The three-category framework of ionic, covalent, and metallic bonding, and the electron behaviour underlying each, is well-established, foundational chemistry, not in scientific dispute. What's worth noting is that real bonds often sit somewhere along a continuum rather than falling neatly into one pure category — many bonds have partial ionic and partial covalent character depending on how unevenly the bonding atoms share electrons, and precisely quantifying that mixed character for a specific bond, rather than treating the three categories as perfectly discrete boxes, is a more nuanced and ongoing area of chemical bonding theory than the introductory three-way split suggests.
This sits inside Chemical Bonding (Ionic, Covalent, Metallic), one of eight topics in General Chemistry, one of six domains in Chemistry, one of seventeen subjects the app can quiz you on.