Chemical bonding and why substances behave differently

Ionic bonding transfers electrons. A metal loses electrons to become a positive ion, a non-metal gains them to become negative, and the oppositely charged ions attract in a repeating three-dimensional lattice. Sodium chloride is the standard example.

Covalent bonding shares electrons between non-metal atoms, producing either small discrete molecules like water and carbon dioxide or giant covalent networks like diamond and silicon dioxide. Metallic bonding is different again: metal atoms release their outer electrons into a shared sea, and the resulting lattice of positive ions is held together by that delocalised cloud.

• Ionic solids melt high, are brittle, and conduct only when molten or dissolved, because the ions must be free to move.
• Simple covalent substances melt low, because melting only overcomes weak forces between molecules, not the strong bonds inside them.
• Giant covalent networks melt extremely high and are usually hard and non-conducting, since every atom is locked into strong bonds. Graphite is the exception, conducting because one electron per carbon is delocalised between layers.
• Metals conduct heat and electricity, and bend rather than shatter, because layers of ions can slide while the electron sea keeps holding them together.

When two bonded atoms differ in electronegativity, the shared electrons sit closer to one of them and the bond is polar. Whether the whole molecule is polar depends on shape as well: carbon dioxide has two polar bonds but is linear, so they cancel, while water is bent, so they do not.

Polarity drives solubility. Polar solvents dissolve polar and ionic solutes, non-polar solvents dissolve non-polar solutes. This single rule explains why salt dissolves in water, why oil does not, and why soap, with a polar head and a non-polar tail, bridges the two.

London dispersion forces exist in everything and arise from momentary uneven electron distribution. They strengthen with more electrons, which is why larger molecules boil at higher temperatures. Dipole-dipole forces act between permanently polar molecules.

Hydrogen bonding is the strongest of these, occurring when hydrogen is bonded to nitrogen, oxygen or fluorine. It is why water has an unusually high boiling point for its size, why ice floats, and why DNA strands hold together yet can be separated for copying.

Questions about properties almost always want the same chain: identify the structure, identify the forces that must be overcome, then state the consequence. Melting a simple covalent solid breaks intermolecular forces only. Conducting requires mobile charged particles, either free electrons or free ions.

The most common error is claiming that melting a molecular substance breaks covalent bonds. It does not, and naming the force explicitly is usually where the mark sits.