This is a vapor pressure experiment, the jar was sealed while the water was boiling on a hot plate, so the steam displaced all the air out of the jar. After sealing, the jar was quickly removed from the heat and let cool. now, there is only water and water vapor in the jar (give or take, there is some hair and dust in the water, and probably some remaining dissolved gases). The pressure inside is always equal to the vapor pressure of the water, which depends on temperature. At room temperature, it's almost a vacuum inside, and as the temperature increases, so does the pressure. At 100C, the pressure inside is one atmosphere, which is the same condition under which the jar was sealed. Whenever there is a heat transfer from the bulk water to the lid (condensing the vapor), the water boils, no matter the absolute temperature. In this case, the water was luke warm, and I could hold the jar while it boiled inside. Another interesting property is how easily the water cavitates. simply moving the jar around or tapping lightly on the top causes the water to separate from the walls of the jar, and it makes a clapping sound.
This is a jar full of only water (liquid and vapor). It boils at any temperature when you apply something cold enough to the top, like ice.
Exactly. it was bottled at atmospheric pressure while it was boiling, so 1 atm and 100 degrees C. Check this graph to see the relationship between the water’s temperature and it’s pressure in the jar (since there is no air, only water vapor). If the vapor is condensed, then the pressure drops below the curve on the graph, that is, the pressure in the jar is lowered below the vapor pressure of the water. Any time the pressure is below the vapor pressure, the water will boil, releasing vapor, until the pressure is equal to the vapor pressure. The pressure does not become negative, it is still positive, just lower than the vapor pressure at the given temperature. You can get below the vapor pressure curve by changing the temperature too, which is what we usually do when boiling water at a pressure near 1 atm (760mmHg)
a slight aside, there is an important difference between the total pressure of the air, and the partial pressure of water vapor in the air. Inside the jar, the two are equal, but in a dry location (not humid) the partial pressure of water vapor is usually less than the vapor pressure of water at that temperature, but since the total large pressure of the atmosphere would not allow a pocket/bubble of very low pressure water vapor to form inside the bulk water, the water cannot boil, but it will evaporate at the surface anyway until the partial pressure of water is equal to the vapor pressure (very humid).
Paint it as a chemical reaction in order to understand its equilibrium state. We basically have:
H2O (gas) ⇌ H2O (liquid)
By sealing the jar with the water already boiling, we initialize the system to be in a state with equal(ish) amount of both liquid and gas. Then we allow the system to cool down so that the liquid water is no longer boiling. Now the system sits at an equilibrium between liquid and gas states.
Now, when we put ice on top of the jar, the water vapor condenses and gets converted to liquid, pushing the equilibrium to the right. But this decreases the overall pressure in the system since fewer particles now occupy the volume above the liquid’s surface. This is essentially the system trying to pull itself back towards the original equilibrium i.e. towards the left of the equation, which it does by making more water vapor i.e. boiling.
This reaction-like picture helps in visualizing the system better in some cases, so I tried to add it alongside the pressure dynamics scenario. You may be interested in Le Chatelier’s principle if you prefer this.
Exactly. it was bottled at atmospheric pressure while it was boiling, so 1 atm and 100 degrees C. Check this graph to see the relationship between the water’s temperature and it’s pressure in the jar (since there is no air, only water vapor). If the vapor is condensed, then the pressure drops below the curve on the graph, that is, the pressure in the jar is lowered below the vapor pressure of the water. Any time the pressure is below the vapor pressure, the water will boil, releasing vapor, until the pressure is equal to the vapor pressure. The pressure does not become negative, it is still positive, just lower than the vapor pressure at the given temperature. You can get below the vapor pressure curve by changing the temperature too, which is what we usually do when boiling water at a pressure near 1 atm (760mmHg)
http://hyperphysics.phy-astr.gsu.edu/hbase/Kinetic/watvap.html#c2
(1 atmosphere is ~760mmHg)
a slight aside, there is an important difference between the total pressure of the air, and the partial pressure of water vapor in the air. Inside the jar, the two are equal, but in a dry location (not humid) the partial pressure of water vapor is usually less than the vapor pressure of water at that temperature, but since the total large pressure of the atmosphere would not allow a pocket/bubble of very low pressure water vapor to form inside the bulk water, the water cannot boil, but it will evaporate at the surface anyway until the partial pressure of water is equal to the vapor pressure (very humid).
You are sollidly correct, and your arguments are correct, but TMI (Too Much Information) applies (to me). Be well, indeed live long and prosper…
Paint it as a chemical reaction in order to understand its equilibrium state. We basically have:
By sealing the jar with the water already boiling, we initialize the system to be in a state with equal(ish) amount of both liquid and gas. Then we allow the system to cool down so that the liquid water is no longer boiling. Now the system sits at an equilibrium between liquid and gas states.
Now, when we put ice on top of the jar, the water vapor condenses and gets converted to liquid, pushing the equilibrium to the right. But this decreases the overall pressure in the system since fewer particles now occupy the volume above the liquid’s surface. This is essentially the system trying to pull itself back towards the original equilibrium i.e. towards the left of the equation, which it does by making more water vapor i.e. boiling.
This reaction-like picture helps in visualizing the system better in some cases, so I tried to add it alongside the pressure dynamics scenario. You may be interested in Le Chatelier’s principle if you prefer this.
Anything for posterity