The discussion centers on the relationship between external pressure and vapor pressure in liquids, specifically addressing how hydrostatic pressure can increase vapor pressure. Participants reference the Poynting correction and Gibbs free energy principles, concluding that while increased external pressure typically suggests greater difficulty for liquid molecules to escape into the gas phase, it actually raises the vapor pressure due to changes in free energy. This phenomenon occurs because the addition of an inert gas increases the total pressure, which in turn affects the equilibrium between liquid and gas phases.
PREREQUISITESChemistry students, thermodynamics enthusiasts, and professionals in chemical engineering or physical chemistry who seek to deepen their understanding of vapor pressure dynamics under varying pressure conditions.
The boiling point is defined as the temperature where the vapour pressure is equal to external pressure. So it will always rise with pressure.realitysickness said:Hi DrDu, do you mind responding to my earlier question: "My only question remains then: given the information you just stated, for situations where temperature is not a constant, why are the boiling point temperatures of a liquid higher at higher pressures?
No, it does not. Ideal gases don't know of each other. So when in the gas phase, the energy of the vapour will only depend on temperature, but not on pressure. The energy of the liquid on the other hand is slightly increased due to compression. So the energy a molecule needs to go from the liquid to the gas phase decreases. For real gases there are slight deviations from ideality at high pressures, but these are weaker than the energy change due to pressure of the liquid.OldYat47 said:Increasing the pressure in a gaseous atmosphere above a liquid increases the amount of energy needed for individual molecules to escape to a vapor state. The first effect is much smaller than the second.