Charles' Law and exploding hairspray cans, a question

[jbd]

getting there. Actually you are there!

There are not fewer molecules. The number of molecules stays the same. They are just farther apart.

Confusion over the names Charles Law and Gay Lussac's Law. In this case theya re the same. Gay-Lussac published his work apparently in 1802. In his report he mentioned the work of Jaques Charles in 1787. Charles invented the hydrogen balloon.

Gay-Lussac also did other work on gases and the combining of gases, for which he produced another law which I mentioned in an early post.

 

[Walter]

Well, I'll be dipped. It appears I was wrong. Gay Lussac's Law is incorrectly referenced in the NOAA Dive Manual. On page 2-11 they refer to Charles' Law as a volume/temperature relationship and Gay Lussac's Law as a pressure/temperature relationship. The NAUI Nitrox manual does the same thing on page 2-6. In the US Navy Manual it refers to volume/temperature relationships and to pressure/temperature relationships as Charles'/Gay Lussac's Law. No mention of Amonton's Law is made in any of these references.

If you'll look here http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch4/gaslaws3.html#boyle you'll be able to read a little about each of the gas laws, including Amonton's Law.

Thanks to jbd for pointing this out! I love learning new things.

One more question for the French speaker's out there. How is "Amonton" pronounced?

 

[Zept]

Originally posted by Ber Rabbit
The balloon rises because heat is applied to the air causing it to expand (the molecules are moving faster and farther apart)
The expanding air inflates the balloon
Continuing to apply heat makes the molecules move even faster (are there now fewer molecules? Is that why the density decreases?)

One difference between a hot air balloon and a scuba tank is that the balloon is not sealed. Air can escape if necessary, so hot air balloons don't explode (they catch fire sometimes, which is equally nasty). Don't think you need to get into that kind of detail, though.

If you want a hands-on demonstration, you could put an empty drinks bottle in the freezer, with the lid off. It will fill up with really cold air. Put the lid on quickly and firmly and take it out. As the bottle warms up, the air inside it will expand. You should be able to feel the difference, and when you loosen the lid, you'll hear the hiss of escaping air. (In Singapore we notice this even with bottles of water from the fridge.)

Zept

 

[Ber Rabbit]

Thanks for the link Walter, I'll have to check it out when I get home from work in the morning! Your version of NOAA must be newer than mine (well, my husbands) mine is January 1975 and the gas laws are in chapter 1. Page 2-11 has the direct effects of pressure on the body (Did I mention I graduated from kindergarten the same year my husband graduated from high school? )

Thanks for the encouragement jbd, I have a b.a. in geology instead of a b.s. because I couldn't pass the physics even with a great tutor. Actually I couldn't even draw the pictures that would help me solve the problems My tutor would say, "that's not the picture they are telling you to draw" and I would say "that's the picture they are telling ME to draw" and he would say "No wonder you're having trouble solving the problem!" The physics test on the instructor exam is going to be a MASSIVE challenge for me!

Zept, I like your demonstration idea, I'm going to have to see if I can figure out how to transport a frozen empty bottle and keep it cold enough for the demo to work. Guess I could always have them go home and do it as an assignment

Thanks!
Ber

 

[Zept]

Originally posted by Ber Rabbit
Zept, I like your demonstration idea, I'm going to have to see if I can figure out how to transport a frozen empty bottle and keep it cold enough for the demo to work. Guess I could always have them go home and do it as an assignment

Coolbox full of ice? As long as you chill the bottle before you seal it, you should be able to leave it sealed for a while. Probably worth a practice run to check the seal. Making your students do the work is always a good idea, though!

Zept

 

[Drew Sailbum]

There is another time that pressure changes brought on by temperature changes comes into play. Try jumping into a :cold: cold spring on a hot summer day. The pressure in your tank quickly drops a bit as the air and tank cools rapidly.

This is easy enough to demonstrate with a pony bottle and an ice chest full of ice water. Works best if you have a pressure gauge with a digital readout.

 

[Ber Rabbit]

The only difference is the way I worded it. I was going from room temperature to hot and you are going from hot to cool. Say your shop hot-filled your AL80 to 3,000psi. You checked the pressure as soon as you were handed the warm tank. Your pressure gauge said 3,000psi. You get to the dive site, hook up your gear and now you only have 2,700psi. If you had waited in the shop for the tank to cool to room temp you would have known you had a short fill and could have asked them to top it off. Some places will overfill a little (or a lot) to compensate for this. I received a fill at a Florida spring, it was a VERY hot fill in August by a man talking about how he wants to dive the Andrea Doria and did I mention he was missing parts of the hand he used to turn the tank valves on and off. I checked my fill and had 3,700psi in my AL80! (was that how he lost part of his hand?) That tankwent straight to the water to cool down! I wasn't going to leave it baking in the Florida August heat possibly causing the air molecules to move even faster and increase the pressure in the tank even more and probably rupture the burst disk! These are the situations we were always taught to relate Charles' Law to scuba diving. That's why I was curious if the exploding hairspray can was the same thing. What we've discovered is that we really need to sit down with the gas laws to figure out if anyone other than myself is confused.

Not looking forward to this physics test, maybe that will be last!
Ber

 

[Walter]

Actually, that's the lesson you should teach. I'd just mention that it is not Charles' Law.

 

[jbd]

I agree with Walter. The point of the lesson regardless of agency is that temperature affects the pressure of a gas when the volume is fixed. You can use Drew's idea to explain what many folks have noticed when they go diving tank pressure drops when you get in the water but have not yet descend. You can also have the students feel a tank that has just finished filling without being in a tub of water. As you know it will be warm to the touch. Youn can then explain that if the tank is left somewhere and exposed to too much heat the pressure will increase to the point the tank will burst. Now you can add that as a means of preventing a tank from bursting there is a burst disc in the valve which will blow out at a pressure well below the bursting point of the tank. In both scenarios the(cooling & heating the tank) Amonton's law is in play.

I'm curious as to why they are drivng away from the scenario presented by Drew Sailbum?

 

[jrtonkin]

Rather than remember a bunch of different laws, and what the individual names of each of them are, I always preferred to remember a single formula:

The Ideal Gas Law

PV=nRT

P is pressure
V is volume
n is the amount of gas (by weight, or by # of molecules)
R is a constant (whose value I can't remember at the moment)
T is temperature

Charle's, Dalton's, Guy-Lussac's, etc. laws are all special cases of this, with one or more of the variables fixed.

So if you double the pressure, and keep the temperature fixed, then the volume is halved; if you double the volume, the pressure is halved. If
the pressure is doubled, but the volume and temperature remain the same, then there's twice as many molecules present.

If you double the temperature, and keep mass and either pressure or volume fixed, the other one doubles. Note that temperature is measured from Absolute Zero (-273.15 deg C, or approx -450ish deg F)

*******

As for exploding aerosol cans:

As the can heats, the amount of material stays constant (aside from some small leaks), and the volume remains near constant (a heated metal will expand, but not a lot), so the pressure goes up.

Given that a campfire probably reaches a temperature around 800 deg C (an estimate), the
pressure in the can will be [ (800 + 273) / (273)]
about 3x what it was with the can at room-temperature.

As the can heats, the metal also gets weaker. This is actually more important than the increase of pressure in making the entire system fail. A can that was 1/2 used up, and then tossed in the fire would "only" be at about 1.5 times the pressure it was at when it was sitting on the shelf in the store. This is probably much less than the breaking pressure of the can. However, because the can has also been heated, the metal is softer, and thus weaker.

For flamability, there are three concerns: the flammability of the contents, the flashpoint of the contents, and the presence of oxygen.

If the contents (either product, or the propellant) are inflammable then you could get an
explosion, rather than "just" a burst can shooting around like an uncontrolled rocket.

If the temperature within the can exceeds the
flashpoint of the contents, and there is 02 within the can (unlikely, since few things use air as a propellant these days, partly for just this reason), then the contents will begin burning within the can. This can dramatically raise the
pressure within the can, both by contributing heat, and by increasing the amount of gas-molecules within it. I.e. T goes up, as does n.

If the flashpoint isin't reached, or there's not enough O2 in the can to support a flame, then the
contents can't burn until the can bursts. Once the can bursts, there's plenty of 02 (from the air) for the combustion, and a direct fire is generally better at causing ignition than just plain heat, so it comes back to the question of whether or not the contents are inflammable.

Ok, enough being an engineer for today, back to pleasanter topics

Jamie