My chemistry teacher once explained it to me like below. Does anyone know how much truth there is to this explanation?
Temperature as measured by a thermometer or your finger is an average. Not every single molecule has the same temperature. The molecules constantly bounce around, smashing into each other, transferring heat to each other. By chance, some molecules will get hit in just the right way by other molecules to reach a very high temperature and then it evaporates. So there is constantly a gradient of temperatures among the molecules and the ones with the highest temperature are the ones evaporating, until there is no liquid left at all.
As the average temperature increases, the chance of some molecules reaching a high enough temperature also increases, so warm water evaporates faster than cold water.
This also explains why evaporation cools down (like when you sweat): the molecules with the highest temperature are the ones evaporating, so the average temperature decreases as those high-temperature molecules leave the system. Only the relatively colder molecules are left behind - thus it cools as a whole.
This also explains why evaporation cools down (like when you sweat): the molecules with the highest temperature are the ones evaporating, so the average temperature decreases as those high-temperature molecules leave the system. Only the relatively colder molecules are left behind - thus it cools as a whole.
The main principle at work here is the enthalpy of vaporization. When matter changes state, there is an associated amount of energy that is absorbed or released - in the case of vaporization, energy must be absorbed. So when sweat forms on your skin and evaporates, it absorbs heat energy from your body in order to undergo that state change.
For water, the energy involved here is remarkably high, much higher than the energy stored by a few degrees difference in temperature. For example, if you wanted to boil off 1kg of water, it would take about 300 kJ to bring the temperature up to boiling from room temperature and over 2000 kJ to boil it all into steam.
There’s a bit more to it, but it’s because of this effect.
There is actually a balance between liquid and gas state, just overwhelmingly in favor of liquid when at normal temperatures. There is a ratio of molecules that will hit each other and transition to gas, and an equal amount gas hitting liquid and condensing. At least when there is a balance between the two sides, aka 100% moisture in the air. Which is not how it is most places.
Normally there is always evaporated water in the air, and anything that evaporated will be moved away in any mildy ventilated area, as you say, it leaves the system. So it never reaches a balance, which is why things dry up at lower temps as water will always evaporate and leave the system.
My chemistry teacher once explained it to me like below. Does anyone know how much truth there is to this explanation?
Temperature as measured by a thermometer or your finger is an average. Not every single molecule has the same temperature. The molecules constantly bounce around, smashing into each other, transferring heat to each other. By chance, some molecules will get hit in just the right way by other molecules to reach a very high temperature and then it evaporates. So there is constantly a gradient of temperatures among the molecules and the ones with the highest temperature are the ones evaporating, until there is no liquid left at all.
As the average temperature increases, the chance of some molecules reaching a high enough temperature also increases, so warm water evaporates faster than cold water.
This also explains why evaporation cools down (like when you sweat): the molecules with the highest temperature are the ones evaporating, so the average temperature decreases as those high-temperature molecules leave the system. Only the relatively colder molecules are left behind - thus it cools as a whole.
@SorteKanin
The main principle at work here is the enthalpy of vaporization. When matter changes state, there is an associated amount of energy that is absorbed or released - in the case of vaporization, energy must be absorbed. So when sweat forms on your skin and evaporates, it absorbs heat energy from your body in order to undergo that state change.
For water, the energy involved here is remarkably high, much higher than the energy stored by a few degrees difference in temperature. For example, if you wanted to boil off 1kg of water, it would take about 300 kJ to bring the temperature up to boiling from room temperature and over 2000 kJ to boil it all into steam.
There’s a bit more to it, but it’s because of this effect.
There is actually a balance between liquid and gas state, just overwhelmingly in favor of liquid when at normal temperatures. There is a ratio of molecules that will hit each other and transition to gas, and an equal amount gas hitting liquid and condensing. At least when there is a balance between the two sides, aka 100% moisture in the air. Which is not how it is most places.
Normally there is always evaporated water in the air, and anything that evaporated will be moved away in any mildy ventilated area, as you say, it leaves the system. So it never reaches a balance, which is why things dry up at lower temps as water will always evaporate and leave the system.