When we increase the atomospheric pressure pressure above the solution , the boiling point of a solution increases. Why does this happen?

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3 Answers 3


Definitions and facts

Boiling point - when the vapour pressure of the solution is equal to the atompheric pressure, then it is said to be boiling.

Vapour pressure - sometimes molecules escape the solution to convert into the gas phase, also molecules in the gas phase combine to turn back into liquid phase. Hence there exists a equilibrium between gas phase and liquid phase.

check this image of Wikipedia

The pressure exerted by this gas phase is called vapour pressure, and the vapour pressure increases with temperature.

For a more accurate definition checkout vapour pressure on Wikipedia and Chemistry.se for an intuitive explanation of boiling points.

Conclusions and interpretations.

So if you increase atmospheric pressure, the solution needs to exert greater amount of vapour pressure to boil.

As mentioned earlier, you need to heat the liquid further so that the vapour pressure matches with the atmospheric pressure. Hence the boiling point increases.

Sources: Wikipedia, NCERT Chemistry textbook for class 12.

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  • $\begingroup$ Sorry, I meant answer. All of the above applied still. $\endgroup$ Commented Apr 13, 2019 at 17:04

I would take a look at this thread A boiling point question.

A crude answer to this is that it takes work for a molecule to "squeeze" or "jump" from one phase into another. Liquids are tightly packed but vapor is quickly moving.

A possible analogy is that moving into the liquid is like squeezing a marble into a full jar of marbles whereas moving into the vapor phase is like trying to run through spaced out but fast moving traffic. Neither is necessarily easier than the other, and actually at equilibrium both require the same amount of effort.

As pressure increases the liquid gets even more packed, and in the vapor the traffic gets even faster moving.

What does this have to do with the boiling point?

Boiling is the process in which molecules move from the liquid into the vapor phase. When the pressure is higher it is harder to move into the vapor. Thus, more energy is required. We get this energy by transferring energy to the liquid in the form of heat. This heat increases the temperature. Hence, to reach a boil at a higher pressure, we need a higher temperature. Temperature is really a measure of how energetically molecules are moving. As pressure increases they need more energy (temperature) to move about and jump into the vapor (traffic).

  • $\begingroup$ Why we refer temperature as temperature of the system but we refer to pressure as pressure of surroundings? Shouldn't both temperature and pressure be properties of the system? $\endgroup$
    – ado sar
    Commented Oct 20, 2020 at 15:58

Suppose there is equilibrium between molecules leaving the liquid (to become gas) and condensing from the gas (to become liquid). When a gas, the molecules occupy a far greater volume than they do when a liquid. Therefore, increasing the pressure displaces the equilibrium (c.f. Le Chatelier) towards liquid, and causes condensation. This would be true at any temperature, but at the boiling point this condensation requires increased temperature to resume boiling.


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