Let's say I have an open container of liquid. According to what I have read, to get the liquid to boil, the vapour pressure has to equal the atmospheric pressure (or there about). I can understand that the pressure the liquid exerts on the atmosphere must approximately equal the atmospheric pressure in order to escape (I feel like it should slightly exceed). However, I am confused as to why it doesn't equal below boiling point.
Let's say I have my particles in the gas phase and not the liquid phase. If the the external pressure is greater, I can force the gas phase into the liquid phase, since it can't "push back" on equal terms. If I keep applying pressure I will continue to compress the liquid until the opposing force of the liquid is equal to the external pressure, and I have reached equilibrium. If the atmospheric pressure was greater than the pressure applied by the liquid, then the liquid would not "push back" enough, and continue to be compressed. It seems to me there must be a point of stability, and this would happen when the pressures are equal. However, this contradicts what happens when boiling occurs. Therefore, does boiling not require that the vapour pressure exceed external pressure. In my case above, the liquid's vapour pressure equals atmospheric pressure at equilibrium, and by increasing the liquid's vapour pressure above atmospheric pressure, the liquid starts to boil.
I realise that this is wrong, but I am not sure why. I don't understand how the external pressure applied to the liquid can be greater the the reverse - the pressure the liquid applies to the surroundings -, otherwise the liquid would continue to be compressed.