36
$\begingroup$

As the title says, I'm interested in knowing if there is any substance — or combination of substances — that ignites (or even increases its chance of spontaneous ignition) when cooled.

I've never heard of such a thing, nor can I find it in Atkins' Physical Chemistry, but I might easily have overlooked it, since I don't know what to call it. Googling gives me information on substances that lower the freezing point or ignition point of explosives/fuels.

It seems that this should obviously be forbidden on thermodynamic grounds, but I can't quite rule out a phase change allowing such behaviour, for instance.

$\endgroup$
  • 5
    $\begingroup$ Ignition is about kinetics, not thermodynamics. You can't cheat thermodynamics, but there may be a chance with kinetics. True, most reactions slow down upon cooling, but this is not an absolute law. $\endgroup$ – Ivan Neretin Jul 25 '16 at 6:26
  • $\begingroup$ I think a lot of substances would ignite if you cool them down from a temperature initially so hot that molecular bonds can't form, but I'm not entirely sure the process would be classified as ignition. $\endgroup$ – user2357112 supports Monica Jul 25 '16 at 20:36
38
$\begingroup$

Actually... yes! Iron(II) oxide is thermodynamically unstable below $848~\mathrm K$. As it cools down to room temperature (it has to do it slowly) it disproportionates to iron(II,III) oxide and iron:

$$ \ce{4FeO -> Fe + Fe3O4}. $$

The iron is in a form of a fine powder, which is pyrophoric (it may catch a fire when exposed to air). You can see it in action here.

$\endgroup$
  • 5
    $\begingroup$ For those of us coming from the hot questions list, 848 K is 574.85 C or 1066.73 F. =) $\endgroup$ – Sidney Jul 25 '16 at 18:51
  • 3
    $\begingroup$ When actually cooled all the way down to room temperature and left alone, $\ce{FeO}$ will stay like that till kingdom come. I guess the disproportionation is possible in a relatively narrow range of temperatures. $\endgroup$ – Ivan Neretin Jul 25 '16 at 21:27
  • $\begingroup$ @IvanNeretin That's true, you have to cool it slowly. Like all chemical reactions, the rate of the disproportionation drops at lower temperatures. $\endgroup$ – vapid Jul 26 '16 at 7:01
11
$\begingroup$

I doubt ignition could ever be favored by cooling since, as for any combustion you need heat. Also the compound must emit enough vapors for ignition to take place and cooling, of course, goes in the opposite direction.

Some explosives, however, could well detonate (not ignite) if the temperature causes them to crystallize. Diazonium chloride salts, for example, are usually not isolated in their crystalline and dry form because they can cause explosions in such a state, while they are safe in aqueous solution.

$\endgroup$
2
$\begingroup$

Phosphorus vapor

Red phosphorus heated at the atmospheric pressure up to approx. $\pu{400 °C}$ sublimates, and upon cooling it builds up white phosphorus:

$$\ce{P_n ->[Δ] n/4 P4}$$

which spontaneously ignites in air at the room temperature being deposited as a film on the surface (bulk sample ignites at $\pu{50 °C}$):

$$\ce{P4 + 5 O2 → P4O10}$$

which is sometimes observed in a form of glowing light (chemiluminescence), as seen, for example, on the famous painting "The Alchemist Discovering Phosphorus" by J. Wright:

enter image description here

A published procedure for the reference [1]:

PREPARATION
Obtain two dry Pyrex test tubes, one 6-inch and one 8-inch. To the 8-inch test tube add 0.25 g red phosphorus. Fill the 6-inch test tube 1/2 full of cold water, dry the outside, and insert it into the larger tube. The hot tip of the smaller test tube will be supported by the neck of the larger tube.

DEMONSTRATION
Heat the red phosphorus in the larger tube until a deposit appears on the cold surface of the inner tube. Allow to cool. Upon removal of the smaller tube the white phosphorus on the bottom mill ignite.


Plutonium

Bulk plutonium ignites only above $\pu{400 °C}$. However, when cooled down, it undergoes a series of phase transitions and its density is increased by approx. $11\%$:

enter image description here

This causes a large hazard issue as plutonium sample cracks while cooled and develops a large surface that swiftly reacts with traces of moist in the air forming plutonium hydrides $\ce{PuH_{2...3}}$ and plutonium(III) oxide $\ce{Pu2O3}$ which are both pyrophoric and ignite spontaneously at ambient temperatures.

References

  1. Brodkin, J. Preparation of White Phosphorus from Red Phosphorus. Journal of Chemical Education 1960, 37 (2), A93. https://doi.org/10.1021/ed037pA93.1.
$\endgroup$
-3
$\begingroup$

It depends on how you cool it. For example, if you try to cool an alkali metal by immersion in water, you can get a rapid exothermic reaction.

$\endgroup$
  • 6
    $\begingroup$ But that's more than just cooling, it's a completely new reagent. $\endgroup$ – fafl Jul 26 '16 at 8:57
  • $\begingroup$ @fafl That's true, and it's the new reagent rather than just the temperature drop which causes the reaction. However, immersion in water is one cooling technique that a lot of people are familiar with, and that specific act of cooling the right/wrong substance would generally cause ignition. $\endgroup$ – WBT Jul 26 '16 at 14:02
  • $\begingroup$ There is also the property of some salts to be cryoluminescent, which means they produce light when they are either precipitated upon cooling down a solution or when the filtered compounds crack upon cooling down. In this case, this is also an electronic transition which produces light, but it is reported to slightly warm up in this process as well. So perhaps you can have an excited state high enough to cause a reaction? Would be the only guess I have. $\endgroup$ – Justanotherchemist Oct 1 '17 at 12:47

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service, privacy policy and cookie policy

Not the answer you're looking for? Browse other questions tagged or ask your own question.