Let's assume that the exact boiling point of water is 100 degrees. Now I have a stovetop and a pot. I add the water to the pot, put it on the stovetop and turn the temperature to exactly 100 degrees. Now what will happen? There's no additional energy flow since there's no temperature difference between the stovetop and the pot right, but the water is already at 100 degrees at it's boiling point so it should boil, but it somehow doesn't. Can you explain this to me? Do we need to keep the stovetop above 100° for water to start boiling?
Edit: Since my question is being misinterpreted here's what I mean. Consider a pot on a stovetop filled with liquid water. Put another cup which can float on top of it in the pot and fill it with water as well. Now turn the stove top so the water in the pot starts boiling. The water in the cup though will not boil while water around it does even though they are at the same temperature 100°C. My conclusion is that you need additional heat rather than the one that just keeps the temperature at 100°C to actually make the water boil, which the cup does not get since all the additional heat from the stovetop is being used to transform the water to a gaseous state in the pot. From this then I can conclude that the stovetop must provide additional energy other then the one that keeps the water at 100°C to make it boil, the stovetop therefore must be at a greater temperature then the liquids boiling point (so greater than 100°C, the assumed boiling point)and cant be just at 100°C. Now why then is this the case. Why do we need that extra energy if the water is already at 100°C which is said to be the temperature at which water boils at normal pressure, if then we can reach the boiling point without the water boiling in the example I gave.