I’m not sure it’s exactly steam, but if I use my fingers to sqeeze my lips tightly shut, and blow, as to produce pressure inside my mouth, when I release the air from my mouth, what appears to be steam is released. I can do it without pressing my lips shut with my fingers as well, but the effect isn’t as dramatic.
Why does this happen? Am I the only freak who can do this, or is it fairly common? I have a feeling it has to do with added pressure possibly on the saliva, but then again maybe not. Any guesses?
I believe the increased presure forces moisture in your breath to condense/begin condensing in the mouth. You then open up very gently and a little puff comes out.
You’re not a freak based on this ability, as I and others I know can do it. You might be a freak for other reaons, but not for this.
In a warm building, hold your fist to your mouth like your cold and shiver. Others will comment, “You’re nuts”, upon which you take the cue and blow out the steam, making it look like a winter chill is in the air.
At the risk of posting a loooooooong response (Chronos ) I’ll keep this brief.
The ‘steam’ is just that, if you consider that steam is water vapor that condenses to wee droplets of water. This is most easily seen outside on cold days. Moisture in your breath condenses in the cold air.
How does this happen inside or on warm-ish days?
Enter the Joule-Thompson effect. When gases expand, they do so by releasing some of their internal energy to their surroundings. This makes them colder*. Even to the point that any moisture in them can condense.
This is why cans of compressed gases (hairspray, spray paint, etc.) get cold when you spray them for a while.
This is nitpicky, but hey, the Net is founded on nitpickiness. When a gas expands, the drop in temperature is not due to releasing heat to the surroundings. If a fixed volume of gas expands to a larger volume it’s pressure and temperature will drop even it retains its original heat energy. It’s because the same heat is more dispersed–the molecules are farther apart and not bouncing around against one another so violently, resulting in a lower temperature. There is a law that gives the relationship among a gas’s volume, pressure and temperature, although exactly which one applies here I don’t remember (Boyle’s?).
Boyle’s law: at constant temprature pressure and volume are inversely proportional
Charles’ law: at constant pressure, temperature and volume are directly proportional
These are combined (with a constant for the amount of gas - not allowed to change) to form the “combined gas law” (Hey, nobody said scientists were creative when it comes to nomenclature, to wit: dark matter), or, if taken at state, the Ideal Gas Law. Problem, there are no Ideal gases, just gases that behave like one at reasonable temperature and pressure ranges.
Recall, that as the gas is expanding, it is “overtaking” part of its surroundings. ANY time a gas changes its state (change in V,P or T) there must be an energy “shift”, either into or out of the current state (set of P,V and T). The Joule-Thompson effect is the Real Gas measurement of what the Ideal Gas Law (Charles’ and Boyle’s laws) states. In this case, the volume is increasing, while the pressure is decreasing. The gas is “doing work” in both cases. This energy “shift” is measured as a drop in temperature.
At any rate, sorry to whip out the Thermodynamics on you (J-T effect) when I could have done just as well (perhaps more so) with high school chemistry (Ideal Gas Law).
