Ionic, covalent and metallic get taught as three separate things with three separate rules. They are not. They are three regions of a single continuum, and where a bond lands on it is decided almost entirely by how unequally the two atoms pull on the electrons between them.
Pick two elements below. The lab works out the electronegativity difference, places the pair on that continuum, draws what the electrons actually do, and tells you what the resulting material should be like — then checks the prediction against a real compound.
Two, because a bond is between two atoms — that is what an electronegativity difference is a difference between. To assemble a whole molecule from as many atoms as you like, use the molecule builder further down: pick a central atom, add bonded atoms one at a time, and it names the result and works out its shape.
A covalent bond needs somewhere to point. Electron pairs — bonding pairs and lone pairs alike — repel each other and spread out as far apart as they can get, and that single idea (VSEPR) predicts the shape of most small molecules to within a degree or two. Build one and watch the geometry follow.
Twelve main-group elements, one central atom and as many bonded atoms as its valency allows. That covers every shape built from up to four electron groups, which is most of the molecules worth meeting first. It stops short of the expanded-octet cases — PCl₅, SF₆, XeF₄ — which need five and six groups and a different set of geometries, and of chains and rings like ethanol or benzene, which need carbon bonded to carbon rather than one centre with ligands.
Water is a bent molecule with a permanent negative end and a positive end, so it can turn its negative oxygen towards a Na⁺ and its positive hydrogens towards a Cl⁻. Enough of those pulls and an ion is worth more in solution than in the crystal, so the lattice comes apart one ion at a time. A hydrocarbon such as petrol or olive oil has almost no charge separation anywhere on it — nothing for an ion to hold onto — so it cannot pay the cost of pulling the lattice apart. "Like dissolves like" is really just this: a solvent has to be able to replace whatever the solute has lost.
Four electron pairs around a central atom spread to 109.5°. Replace a bonding pair with a lone pair and the angle closes slightly — a lone pair is held by only one nucleus, so it spreads wider and squeezes the bonds. That is the whole reason ammonia sits at 107° and water at 104.5° rather than at the textbook 109.5°.
Every reaction is the same trade: pull some bonds apart, which always takes energy in, then let new ones form, which always gives energy out. The sign of the difference is the entire story. Get more back than you spent and the mixture heats itself — that is a fire. Spend more than you get and the reaction will only run while something keeps paying for it.
The same three names, side by side, with what each one actually does to the material you can hold in your hand.
Hit an ionic crystal hard enough to slide one plane of ions over the next and every contact flips from opposite-to-opposite to like-to-like. The layer that was being held is now being pushed, and the crystal splits along that plane. Do the same to a metal and nothing changes: the cations are all positive already, and the electron sea that binds them does not care where they are. So the metal deforms instead of cleaving — that is malleability, and it is the same electron sea that carries the current.
Covalent bonds hold a molecule together; something much weaker holds one molecule to the next, and that is what sets melting and boiling points. Follow the group 16 hydrides down and the boiling point rises steadily as the molecules get bigger — except water, which is 180 °C out of line. Hydrogen bonding between its very positive H and its very negative O is worth roughly 20 kJ/mol per bond, about a twentieth of a real covalent bond and quite enough to keep the oceans liquid.