Water is said to "boil" when bubbles of water vapor grow without bound, bursting at the surface. For a vapor bubble to expand, the temperature must be high enough that the vapor pressure exceeds the ambient pressure (the atmospheric pressure, primarily). Below that temperature, a water vapor bubble will shrink and vanish.
Superheating is an exception to this simple rule; a liquid is sometimes observed not to boil even though its vapor pressure does exceed the ambient pressure. The cause is an additional force, the surface tension, which suppresses the growth of bubbles.
Surface tension makes the bubble act like an elastic balloon. The pressure inside is raised slightly by the "skin" attempting to contract. For the bubble to expand, the temperature must be raised slightly above the boiling point to generate enough vapor pressure to overcome both surface tension and ambient pressure.
What makes superheating so explosive is that a larger bubble is easier to inflate than a small one; just as when blowing up a balloon, the hardest part is getting started. It turns out the excess pressure  due to surface tension is inversely proportional to the diameter  of the bubble. That is, . 
This can be derived by imagining a plane cutting a bubble into two halves. Each half is pulled towards the middle with a surface tension force , which must be balanced by the force from excess pressure . So we obtain , which simplifies to .
This means if the largest bubbles in a container are small, only a few micrometres in diameter, overcoming the surface tension may require a large , requiring exceeding the boiling point by several degrees Celsius. Once a bubble does begin to grow, the surface tension pressure decreases, so it expands explosively in a positive feedback loop. In practice, most containers have scratches or other imperfections which trap pockets of air that provide starting bubbles, and impure water containing small particles can also trap air pockets. Only a smooth container of purified liquid can reliably superheat.