Terminator-Proof Your Life With Homemade Liquid Nitrogen

Illustration for article titled Terminator-Proof Your Life With Homemade Liquid Nitrogen

Scifi fans know that debate about gun control is besides the point. Guns won't save anyone from rampaging Terminators, aliens, or Jason Voorhees. Liquid nitrogen will. Here's how to make some on a super-low budget, for home protection!

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As everyone knows, no home is safe without a security company sticker on the door (to stop regular criminals) and a vat of liquid nitrogen in the hall (to take care of today's alien, robot, supernatural, or genetically engineered supercriminals).

But why pay for pricey liquid nitrogen delivery (or bother the neighbors for a cup of nitrogen when you run short) when you can just make it yourself? Follow these easy steps to making some liquid nitrogen. It is so much easier than just giving up your paranoia.

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Now, this is a secret that Big Nitrogen doesn't want you to know: Nitrogen is all around you. Yes, even now! What you're breathing in and out is almost eighty percent nitrogen. You're living off the twenty percent oxygen in air (and probably being slowly killed by the trace amounts of other elements).

Don't you feel silly? Here it was all around you, and you were paying for it. But don't feel bad. Just get started!

The first step in the process is compressing the regular old air you get for free. You'll need to get the pressure up to 3,000 pounds per square inch, so I suggest everyone in your household piling on the inner tube you got for the pool.

For the next step, you will need that perfect metaphor for the Internet: a series of tubes. This particular series of tubes will be a way for the gas to shake of that heat from compression, the way a lunatic shakes off being crazy if you shove them in a straight jacket. Keep it pressurized, keep allowing it to radiate heat. Soon it will become cooler.

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Like passengers on an over-crowded bus, a recently-released gas will begin to expand. Unlike the passengers, however, it won't do it via the judicious use of elbows. Gases need energy to expand, and they get energy from their environment. That's why they feel cool. They "suck" the heat out of their surroundings in order to get that energy.

The kids will most likely complain that they're getting cold. Just tell them that if they don't take their turn helping out, they will have to deal with it when Pumpkinhead comes a-knocking.

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Illustration for article titled Terminator-Proof Your Life With Homemade Liquid Nitrogen

That will shut them up.

Finally, the gas will get so cold that, even when the pressure is lifted, it won't immediately expand. It will be compressed into liquid form.

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The rest is simple; pour it quickly and neatly into your vacuum-insulated thermos. Do try to save up and buy your own. The rented ones are a second-rate and expensive. Et voila!

Illustration for article titled Terminator-Proof Your Life With Homemade Liquid Nitrogen
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Next time: How to make sure a vat of boiling iron swallows mechanical pursuers and not your pets. Fences aren't enough!

Via: How to Do Things and Ask the Van.

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DISCUSSION

I'm sorry, Esther, but your analogies (and the two links you offered) don't actually explain why or how nitrogen becomes a liquid, nor do they explain how the environment gets colder in a way that's entirely correct.

There is a major step missing between your paragraph that starts with

"The first step in the process is compressing the regular old air you get for free."

and the paragraph that starts with

"Like passengers on an over-crowded bus, the compressed gas will begin to expand."

and the absence of the stuff in between actually suggests something that isn't correct at all (your statement that the gas will begin to expand).

Nitrogen, like most gasses, is very nearly what's known as an ideal gas, that is, a gas that follows the Ideal Gas Law, PV = kT, pressure times volume is proportional to temperature (k is a constant).

If you don't change the size of the container where your nitrogen is, its volume will be constant, so pressure becomes proportional to temperature (and vice-versa). This is because, in the absence for a reason for the gas volume to decrease, the gas will always occupy the entire volume of its container.

So, by increasing the pressure, the gas will heat up. That makes sense intuitively, since higher pressure under a constant volume means more frequent collisions among the molecules, meaning higher molecular speeds, meaning higher average kinetic energy per molecule, which is exactly what temperature is.

Note, however, that the gas, now at high temperature and pressure, will NOT expand (contrary to your statement), because its volume cannot be larger than the volume of the container, which is fixed.

To be fair, your statement refers to the expansion of the gas after the liquid nitrogen is exposed to the environment, and is then correct. That's not what is implied, though, because there is a step missing.

The missing step is that the gas is allowed to cool down by exchanging heat with something external to it. This is still part of the process of making liquid nitrogen.

By allowing the high-presure/high-temperature gas to come to thermodynamic equilibrium with a large external environment at room temperature (a so-called heat sink), the temperature of the nitrogen comes down to room temperature.

In the process, the nitrogen suffers a phase transition from gas to liquid. This also makes sense intuitively, since at high pressure and (relatively) low temperature, the volume has to decrease (ideal gas law once again), bringing the molecules closer together, which allows their normally weak bonding to take effect over large distances, turning the gas into a liquid. Note that the container having a fixed volume does not prevent the gas from contracting. It only prevents the gas from expanding beyond the container's volume.

So, now, after this "missing step", we have a quantity of liquid nitrogen, at high pressure and room temperature. All you need is for your container to be able to withstand the high pressure. It need NOT be thermally insulated.

If, now, you expose some of that liquid nitrogen to an environment of comparatively lower pressure and unconstrained volume, then the exposed liquid nitrogen will evaporate. In doing so, it absorbs heat from the environment, thereby cooling it down.

In more detail, this is what happens.

The liquid nitrogen is at high pressure. The environment is at normal atmospheric pressure. Because of the pressure difference, the molecules inside the liquid will push away the molecules on the surface of the liquid, making them evaporate.

At first, the evaporated molecules are at high pressure but they are now free to expand and they do so, at the expense of their own internal energy. As a result, their temperature drops from their initial value (room temperature). Since, however, they're in contact with the environment (which is still at room temperature), heat starts to flow from the environment to the evaporated molecules (by means of collisions between air molecules and the evaporated nitrogen molecules), fueling their expansion and decreasing the environment's temperature. This is where the second paragraph I quoted comes in.

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Incidentally, I noticed that you like to use the "people crowded in a bus" analogy to explain increase in pressure and expansion, in gasses. You've used it in previous articles.

That is a bad analogy, because there is a profound difference between people in a crowd and molecules in a gas: the people are not moving.

Even if you talk about a crowd of people moving, it's still a bad analogy, because the motion of the crowd is organized; the people in the crowd will be moving along the same direction. In a gas, even in a gas flowing through a pipe, there is random motion along all directions, and that is what accounts for much of the behavior of gasses.