I've been confused about this for a while now and haven't found anything that directly addresses this problem. Please correct any faults you see in my reasoning. My understanding of vapor pressure is that it is the pressure exerted by the vapor phase of a substance above the liquid phase onto the liquid phase (and container walls in a closed system). Molecules are always leaving the liquid phase due to their kinetic energy and entering the vapor phase. In a closed container, these molecules eventually enter an equilibrium state where the number entering the liquid phase equals the number entering the vapor phase. However, in an open container (e.g. pot on a stove) equilibrium cannot be reached due to the vapor being dispersed into the atmosphere which is not the case in a closed container. Now, the definition of boiling is when vapor pressure reaches and/or exceeds the atmospheric pressure. If equilibrium cannot be reached in an open container, how is vapor pressure maintained above the liquid such that boiling can occur? Boiling in open containers occurs all the time so there must be some way that the vapor pressure above the gas can exceed atmospheric pressure even if it's being dispersed otherwise boiling wouldn't occur. Side note: If the vapor is above the gas, does this not make it the atmospheric pressure above the liquid phase. Thus, the Earth's atmospheric pressure is intertwined with the vapor pressure making the atmospheric pressure above the liquid phase now a combination of vapor pressure and the Earth's atmospheric pressure? This is part of where my confusion is coming from because it seems that overcoming atmospheric pressure is a redundancy because vapor pressure is part of the atmosphere above the liquid thus vapor pressure would have to be overcoming itself. My only explanation for this is that the vapor pressure is also within the liquid phase and is exceeding the vapor pressure due to the gaseous phase above. Not sure if this is correct though.