boiling-point elevation
Pure water boils at 100°C, but stir in salt and it must be pushed a little hotter before it boils. Boiling-point elevation is exactly that: dissolving a non-volatile solute raises the temperature at which a liquid boils. The solute, in effect, makes the liquid hold on to its molecules a little more tightly.
The reason traces back to Raoult's law. Solute particles lower the solvent's vapour pressure, and a liquid boils only when its vapour pressure climbs up to match the surrounding air pressure. With vapour pressure pulled down, the liquid must be heated further to reach that boiling point. The size of the rise depends only on how many solute particles are present — it is a colligative property — and is found from the solution's molality times a constant fixed by the solvent.
The effect is real but usually small: ordinary cooking salt raises water's boiling point by only a fraction of a degree, far too little to cook food noticeably faster. It is the same physics, though, that lets engine coolant run hotter without boiling over, and it gives chemists another particle-counting route to a substance's molar mass.
Salting a pot of pasta water does raise its boiling point — but typical amounts add only a few hundredths of a degree, so the salt is for flavour, not for cooking speed.
A dissolved solute makes the solvent boil a little hotter than when pure.
The solute must be non-volatile (it should not evaporate easily itself) for the simple rule to hold. A solute that splits into ions counts as multiple particles, so it raises the boiling point more — by a factor of roughly the number of ions it produces.