![]() Or with microwaved coffee, provide a dry, air-filled wooden stir-stick. To prevent this, use 'boiling stones' sold in laboratory supply catalogs. Sometimes they're violent enough to shatter glass. Instead it superheats far above 100C, then unexpectedly produces a few spontaneous micro-bubbles, and exhibits the boiling-explosions called "bumping." The explosions may splash boiling water out of the container. Often with new glass cookware, (with no scratches,) and with water heated on a gas stove (with no tiny hot-spots,) the water won't boil. But without this large temperature excursion at the metal surface, visible boiling may not commence. When the main volume of water reaches 100C, seed-bubbles are already present on the hot metal, so the pot will immediately begin a visible boil. The metal surface will be covered with spontaneous steam pockets, but these bubbles cannot grow, since they're right against cooler, under-100C water. On a typical stove with a metal pot, the metal bottom will be heated far above 100C, even though the water has not yet approached 100C. A fully-wetted rough surface won't prevent superheating. The hot water fills the microscopic roughness with steam, which then condenses, removing any air pockets that let the rough surface act like a "seed" for roiling boil. But liquids may refuse to boil even when up against a very rough surface, if that surface has been previously wetted with water over 100C. If the surface of your ceramic mug lacks air-filled micro-scratches, the coffee will not boil until its temperature is raised far above 100C. So, yes, your coffee heated in a microwave oven can superheat and explode, even though coffee is very impure water. ("Mythbusters" show got it wrong!) Instead, boiling is seeded by existing micro-bubbles trapped in small crevices. Boiling-bubbles are typically not seeded by dirt or contamination. Without these, all boiling takes place silently, at the surface where water touches air.īeware of common misconceptions. For vapor pressure to exist within water, first gas pockets must exist within water. Instead, when the vapor pressure is equal to the external pressure, then any existing bubbles will begin growing continuously.Īnd, if no bubbles are already present, then the water will superheat far above the boiling temperature, yet no bubbles will appear. When the vapor pressure is equal to the external pressure, there will form a bubble. vapor pressure that equals the external pressureĪre they both causes of the boiling (coming up of the bubbles)? Or is the vapor pressure that equals the external pressure the cause? When the boiling point is reached, there happen two things: In this video they say boiling has more causes. And in my opinion a cause happens BEFORE the consequence. I doubt if it is a link of causes (the one thing causes the other) because they happen at the same time. So, my question is: is this a chain/link of causes? So the first link causes the next one? So the temperature increase causes the evaporation to increase which causes the vapor pressure to increase which causes the forming of a bubble (the actual boiling)? When the evaporation increases, the vapor pressure will increase too. As the temperature of the water increases, the evaporation increases.Ģ. Everyone said something else about the cause of the boiling. I've searched alot about the boiling of water and this confused me. I'm a first year physics student and from the Netherlands. As pressure increases they need more energy (temperature) to move about and jump into the vapor (traffic).I've got a question about the boiling of water. Temperature is really a measure of how energetically molecules are moving. Hence, to reach a boil at a higher pressure, we need a higher temperature. We get this energy by transferring energy to the liquid in the form of heat. When the pressure is higher it is harder to move into the vapor. What does this have to do with the boiling point?īoiling is the process in which molecules move from the liquid into the vapor phase. Neither is necessarily easier than the other, and actually at equilibrium both require the same amount of effort.Īs pressure increases the liquid gets even more packed, and in the vapor the traffic gets even faster moving. ![]() Liquids are tightly packed but vapor is quickly moving.Ī 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. I would take a look at this thread A boiling point question.Ī crude answer to this is that it takes work for a molecule to "squeeze" or "jump" from one phase into another.
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