[QUOTE]
*Originally posted by thinksnow *
**
Guess I must be a Seabee! 
Quasi
Some questions for clarification:
Was it incredibly humid, or very dry? With all the warm beer, I’m assuming it was hot.
Were the drums just open at the one end, with the small one sealed? Or could the smaller drum have openings in the bottom?
Were these steel cans, or aluminum?
Would the SeaBees have access to semi-exotic liquids like ammonia? Rubbing alcohol?
Man. Is this a guy thread or what?
Jill
Danalan:
This is what I remember. It was Vietnam, and it always seemed hot to me. It also always seemed humid to me too, but then again I grew up in Colorado where humidity rarely got above three percent or so.
As I remember they were steel beer cans. I don’t know of any openings at the bottom of the smaller drum, but they could have been there. When I returned with the fuel, the set up was completed and just waiting for me. When they started playing with the aviation fuel, I stayed a distance away.
We had both ammonia (cleaning supplies)and alcohol (medical supplies) around the camp.
TV
Incidentally, the Seabees were founded by Rear Admiral Lewis Coombs, a graduate of Rensselaer Polytechnic Institute, where I’m almost through being a student. Nothing of substance to add, though.
Johnny L.A. wrote:
Even so, it’s not nearly as bad a situation as with a lead-burning automobile engine.
When an engine burns leaded gasoline, the lead comes out as a vapor. If this vapor is released at ground level, it becomes part of the air (temporarily) and everybody has to breathe it. If, however, this vapor is released at altitude, it condenses out into particulate lead by the time it reaches the ground. So instead of having to breathe lead vapor from an airplane engine burning leaded avgas, we only have to deal with a few microscopic grains of lead settling onto the countryside. Except while the plane’s taxiing, of course.
A bigger environmental no-no, which modern private pilots are still encouraged to do, happens when a pilot checks the fuel in the plane’s tanks during a pre-flight inspection. The pilot drains a little bit of fuel out from a tiny fuel-drainage spigot at the bottom of each tank, checks it for water and sludge deposits, and then dumps this little bit of fuel out onto the pavement. I’m waiting for the EPA to go bonkers over this practice one of these days.
Back to the OP…
The technique of cooling beer in that manner is not new. I can remeber being told about it many years ago by ex-army people. The only differnece was that they buried the beer cans in the ground, then poured a ring of petrol around the buried cans. The fuel was then lit, and when it had burned off, the cans were dug up, at a lovely cool temperature, just right for consumption.
They also mentioned the fire extinguisher trick, but fuel is generally much cheaper and more effective (so I was told).
That’s true, although I wouldn’t have put it like that. It sounds like you’re saying “Pilots are encouraged to dump fuel on the ground.”
It is extremely important that the fuel be checked for contaminants before each flight. See what I wrote earlier about needing the engine to remain operating in flight. The fuel must also be checked to make sure it’s the right fuel. Reciprocating aircraft engines don’t run very well on low-octane jet fuel, and gasoline in a jet engine can damage it. So a fuel sample must be taken as part of the preflight.
But what about the technique? When I learned to fly, I was taught to disperse the sample with a swift sweep of my arm. The avgas would evaporate before it hit the ground. On the other hand, the engine sump (there’s a sump in each wing tank and one forward of the firewall in a Cessna) was simply drained onto the ground for a few seconds. The idea was to drain out any water, and to see if there was any water in the fuel puddle on the ground. When I started flying helicopters I collected this sample in the collector so that I could take a close look at it. I never had water in my fuel, so I was never sure if I’d recognize it in a puddle on the ground. Besides, what if it had been raining and water was already there? In any case, now that I collect the fuel sample in the sampler, I can evaporate it before it hits the ground with a sweep of my arm.
There is still the question of pollution, whether pilots “collect and disperse”, drain/pour the fuel directly onto the ground, or dump it in an empty coffee can. (No, you can’t put it back in the tanks. Not recommended for “clean” fuel, and a dumb move for contaminated fuel.) It’s probably not a good idea to carry it in an empty mayonaise jar in the cockpit until you can take it to a hazardous-waste site either. Many airports discourage dumping fuel samples, but there’s not a whole lot they can do.
OK then, back to the OP.
Here’s what I think may be going on:
The fuel is in the open barrel, slowly vaporizing due to the heat. When the vapor is ignited, it oxidizes rapidly. The oxidation doesn’t extend far into the barrel, because there is no available oxygen for the reaction. This leaves plenty of fuel in the barrel. One of the effects of the rapid oxidation above the barrel is a reduction in pressure. This causes the fuel left in the barrel to rapidly vaporize (the vapor pressure directly above the fuel is reduced drastically?), and rise up out of the barrel to be oxidized.
The rapid vaporization of the remaining fuel is what causes the temperature reduction. Just like the evaporative cooling of wet clay pots, except at high speed. I think the second barrel was there mainly to keep the taste of the fuel off the cans.
First off, I would like to debunk the pressure drop theories. There will be only minute pressure gradients associated with an open flame. They can be ignored. The idea that oxygen being consumed creates a pressure drop is even farther off. For every molecule of oxygen consumed, you get either one carbon dioxide molecule or two molecules of water. For a generic hydrocarbon fuel,
(m + n/4) O + C H -> m CO + n/2 H O
2 m n 2 2
Note that we end up with about (n/4) more molecules after combustion. This would tend to increase the pressure. After all, when we do have a large pressure gradient in a flame, the pressure tends to be higher in the combustions volume. Events of this type are also known as “explosions.”
However, setting fire to a fuel will cause it to evaporate more quickly. Obviously, a flame in close proximity would heat the liquid which helps it evaporate, but works against cooling our beer. The second mechanism is that the partial pressure of the fuel in the vapor phase is lower (not the “real,” mechanical pressure) than would otherwise be as the fuel in the vapor phase is consumed by the combustion reaction.
I find it a little hard to believe that beer is cooled by the increased rate of evaporative cooling versus heat gains from the combustion rxn, but… if that is what is observed… Still, you might just be better off putting together your cooling contraption and placing it in front of a big fan over setting fire to the thing. Any of the resident pyromani^H^H^H^H^H^H^H combustion researchers up for some controlled experiments?
What if the beer is held to the bottom of the barrel completely submerged in the fuel and away from the flame? Wouldn’t that negate the flame’s effects on the beer and enhance the cooling gradient?
–Tim
cjclark wrote:
Now, in a monospaced pre-formatted font, for those of us who can’t tell where those numbers and letters on the 2nd line are supposed to go:
(m + n/4) O + C H -> m CO + n/2 H O
2 m n 2 2
The {CODE} directive. It’s your friend. [TM]
Even better is to use subscripts, thus:
(m + n/4 O[sub]2[/sub] + C[sub]m[/sub]H[sub]n[/sub] -> m CO[sub]2[/sub] + n/2 H[sub]2[/sub]O
You get subscripts by using the sub and /sub tags, so if I type
H[[sup][/sup]sub]2[[sup][/sup]/sub]O
I’ll get
H[sub]2[/sub]O
Superscripts are similar, but with sup instead of sub.
I fail to see how a generic formula for combustion (even one with the correct gaseous products) debunks the theory that a pressure drop occurs at the location of the evaporation reaction. I recognized that the oxidation reaction has more gaseous molecules on the product side compared to the reactant side, but for there to be a cooling effect the evaporation and oxidation reactions need to take place in different locations. Thus the number of gaseous molecules on the product side of the oxidation reaction does not directly control the pressure at the location of the evaporation reaction.
Pressure gradients associate with an open flame may be small, but to call a gigantic fireball stretching way up in the sky a mere open flame is a bit of an understatement. With so many molecules being drawn up into the flame so quickly, it certainly seems possibly that a drop in real pressure might occur (in addition to he decrease in the partial pressure of fuel). I have limited fireball experience so I can’t say for sure whether or not a real pressure drop will occur. It seems logical to me that if you if created a big enough fireball, you could get a pressure drop.
(You are correct to clear up any notion that the pressure drop was caused by comsumption of oxygen. I didn’t mean to make that suggestion.)
It does not really need to. If you were to have a combustion reaction that took place on the timescale associated with a combustion reaction and it had the opposite effect, you would have an “implossion.” Macroscopic pressure changes that take place on the timescale of combustion reactions are fairly catastrophic events.
The rising of a fireball, if I recall correctly from wa-ay back in my chemical engineering days, would be associated with the Grashoff number which in turn is a convective effect. Convection is slo-ow. A rising fireball is the result of bouancy (a blob of hot air), a density gradient, not a pressure gradient. The convective transport would have its greatest effect cooling from causing new, cool air with zero fuel concentration to rush in over the flame as opposed to any pressure effects.
I should really check up on this list more often.