# How is the temperature in which vapor pressure is equal to atmospheric pressure the boiling point of the liquid? [duplicate]

I was just learning about vapor pressure and got very confused. Vapor pressure was defined as the pressure exerted by a vapor in thermodynamic equilibrium with its condensed phases at a given temperature in a closed system, which was the pressure the vapor exerts on the closed container when it reaches the dynamic equilibrium. However, this is merely a calculation AFTER it reaches equilibrium, and I don't understand how vapor pressure can also be interpreted as the measure of the tendency of a material to change into the vapor state. Also, if it is the lack of enough vapor pressure, or "tendency of the liquid to change to vapor" compared to the atmospheric pressure that's keeping the liquid from vaporizing, how can evaporation occur with lower pressure than the atmospheric pressure(because if the vapor pressure is higher than atmospheric pressure it is already boiling), since evaporation is supposed to be the molecules on the surface of the liquid randomly gaining enough energy to overcome the atmospheric pressure, implying that it has greater pressure than the atmosphere. I feel like I am overthinking the topic, but I didn't just want to give up thinking about it, so please help if possible!