Dissolving anything in a solvent lowers its freezing point and raises its boiling point. Both effects depend on how many particles are dissolved rather than what they are, which is what makes them colligative properties.
What surprises people is how lopsided the two are.
Freezing: the effect is large
Water has a cryoscopic constant of 1.86 degrees per molal. Dissolve 100 grams of salt in a kilogram of water and the freezing point drops to about minus 6.4 degrees Celsius. Two hundred grams takes it to minus 12.7.
That is a real, useful effect, and it is exactly why salt goes on roads. It does not generate heat and it does not attack the ice chemically. It lowers the temperature at which the water film on the ice can remain solid, and below that temperature the ice simply cannot persist.
The freezing point depression calculator takes the van t Hoff factor as an input because it matters enormously here: sodium chloride splits into two ions and is therefore roughly twice as effective per mole as sugar, which stays whole.
Boiling: the effect is nearly nothing
Water has an ebullioscopic constant of 0.512, less than a third of its cryoscopic constant. Ten grams of salt in a kilogram of water raises the boiling point by 0.175 degrees.
A heavily salted pasta pot might gain half a degree. The claim that salting the water makes it boil hotter and cook faster does not survive the arithmetic: the temperature change is too small to measure with a kitchen thermometer, let alone to affect cooking time.
Salt goes in pasta water for flavour. That is a perfectly good reason, and it is the only one the boiling point elevation calculator will support.
Why the two constants differ
Both constants come from the enthalpy of the relevant phase change and the solvent's properties. Melting and vaporising are different transitions with different energies, so there is no reason for the two constants to match, and for water they differ by a factor of about 3.6.
Other solvents are more extreme still. Camphor has a cryoscopic constant of 37.7, which is why it was historically used to determine molar masses by freezing point depression: a tiny amount of solute produces a temperature change large enough to measure accurately.
Road salt has a floor
Below about minus 10 degrees, salting stops being practical. The quantity needed grows, the simple proportional relationship starts to overstate the real effect as ions begin pairing up, and eventually you are spreading more salt than the road can hold.
That is why colder regions switch to calcium chloride, which dissociates into three particles rather than two, or to sand, which does not melt anything at all and simply provides grip.
For the underlying cause of both effects, see the Raoult's law calculator. For the concentration measure they both depend on, see the molality calculator.