Reusable handwarmers are made of something dissolved nearly to its saturation point in hot water (Sodium Acetate in the case of HotSnapz) at a high temperature, and then cooled to a temperature at which some would normally be insoluble, but remains dissolved due to a lack of a site to grow a crystal from. When the crystal finally grows, triggered by popping a little disk, the process is exothermic and it releases heat.

What I'm wondering is: how does popping the little metal disk inside the hand warmer allow nucleation to begin?

EDIT: I got a bit more information, but nothing definitive. What I’ve heard is that the metal disk may either be initiating nucleation by releasing tiny metal particles when it snaps or by snapping quickly enough to cause cavitation. The first seems unlikely, because if bits of metal were released and functioned as nucleation sites, the warmer wouldn’t be reusable - the crystals could just reform on those metal bits. Cavitation seems more plausible, though I still don’t understand how that would allow crystallization to begin.

EDIT 2: I did some experimenting, and a whack also starts nucleation, not just the disc snapping. So, any shock can cause the sodium an acetate ions to come out of supersaturated solution. How?

  • $\begingroup$ I am sure the thermodynamics of supersaturated solution must be complex. In older times in order to initial crystallization in a beaker, one would slightly rub the beaker with a glass rod. A single scratch or a microscopic (invisible) glass particle would initiate crystallization. Similarly, I think, bending the metal disk provides new nucleation sites on its surface. $\endgroup$
    – AChem
    Sep 16, 2020 at 3:10
  • $\begingroup$ Recall superheated water in a microwave which is not willing to boiling, the moment you add a spoon or something solid it boils violently. It must be a similar phenomenon. $\endgroup$
    – AChem
    Sep 16, 2020 at 3:11
  • $\begingroup$ Related: pnas.org/content/113/48/13618 $\endgroup$
    – Karsten
    Sep 16, 2020 at 3:47
  • $\begingroup$ I don't have first hand experience with super-saturated solutions, but could it be that they precipitation would occur on stirring? The snap could be considered a vigorous quick stir. $\endgroup$
    – TAR86
    Sep 16, 2020 at 5:21

2 Answers 2


Edit: Contribution of enthalpy and bond formation

Edit 2: Changed the description of entropy

After thinking for sometime, I realized that I neglected to mention three things:

  1. The Driving force of precipitation reactions depends on the charge and size of the ions which corresponds directly to the magnitude and order of the hydration shells formed upon dissolution of the salt.

  2. The enthalpy as a driving force of the precipitation of the salt. (Remember the Gibbs Free energy equation has two terms, entropy and enthalpy)

  3. Additionally, the formation of ionic bonds can be lower in energy (more stable) than the electrostatic interactions present in solution. The release of energy during salt formation is transferred to the surroundings as heat. This would be a negative enthalpy (heat flowing out of the system) and if large enough, would overcome any decreases in entropy stemming from the formation of highly order crystalline structures and drive the precipitation.

In this case it is probably the enthalpy term that is driving the reaction, since you are putting a ton of heat into the system to dissolve the salt, once the solution cools, there is a lot of potential energy that is stored in the position of the atoms in the system. This is then released as heat upon the formation of the crystal structures.

Entropy changes in precipitation reactions:

In order to make a supersaturated solution, you have to add energy to form hydration shells around the ions (entropy in the system is decreasing). The thing to remember is that these hydration shells decrease the dispersal of the ions relative to that in the crystalline structure for some salts. When you heat and mix the solution with excess solute, you are providing sufficient energy to form hydration shells. When the solution cools (Slowly), the hydration shells remain intact, however, the system is unstable and applying enough energy to the solution will disrupt the hydration shells and crystallization of the salt will occur. The large amount of heat that is released is coming from the increase in the micro-states available for the ions (gain in entropy) in removing the hydration shells around the ions. The hydration shells are less dispersed than the crystalline structures of the salt and that is why entropy increases during some crystallization reactions, making this process exothermic.

Here is why the disc snapping is important:

Even though this process is very exothermic, it still requires activation energy. The mechanical force you are referring to (i.e. snapping the disc or whacking the heating pad) is doing just that by simply transferring energy to the molecules, through vibrations, which provides the activation energy for crystallization.

Just as walking can cause energy transfers from the mechanical forces of your legs to the molecules of the floor (i.e. through friction) and increase the temperature of the surface, some of the force applied on the disc transfers to the solution and if it is enough, will overcome the activation energy barrier of the process and crystallization will begin.

  • $\begingroup$ Interesting that the supersaturated solution - a liquid - is more ordered than NaOAc.3H2O crystals plus some saturated solution. $\endgroup$ Sep 19, 2020 at 19:48
  • $\begingroup$ @JamesGaidis I thought about this specific case and realized that it is probably the enthalpy that is the driving force since sodium is a non-acidic cation with a relatively small Z^2/r value. However, I would refer you to this link if there is any doubt about entropy as the driving force for other precipitation reactions. people.wou.edu/~courtna/ch412/ppt2.htm $\endgroup$
    – user98623
    Sep 19, 2020 at 21:03
  • $\begingroup$ Interesting link! But entropy doesn't drive, it is an explanation, just like enthalpy is a result. What two forces, or rather what force plus what blockage are operating here? NaOAc .3H2O crystal structure has Na+ surrounded by 6 oxygens in a distorted octahedron (onlinelibrary.wiley.com/doi/abs/10.1107/S0567740876002367). But hydrated Na+ is all over the place, from 4 to 8 waters (ncbi.nlm.nih.gov/pmc/articles/PMC3250073). Cooling a solution must enable a slide into one metastable (yet fluid) structure rather than the more stable crystalline form. $\endgroup$ Sep 20, 2020 at 13:08
  • 1
    $\begingroup$ @M.Farooq I apologize for the inaccuracy of my answer and will be corrected tomorrow morning. I take responsibility for propagating such “garbage”. From my current understanding, Boltzmann would describe it as an increase in micro-state’s of the system. Even though some processes may lead to more disordered systems, there isn’t really a way to quantify it and it is not necessary for disorder to increase to have a positive entropy? $\endgroup$
    – user98623
    Sep 22, 2020 at 1:15
  • 1
    $\begingroup$ @dval98, Entropy is indeed a complex topic and very challenging. I tried to read Clausius original work (translated from German) but it is so engrossed in hardcore heat engines and mechanics that I gave up. One needs a very good grasp of thermodynamics and mathematics in which chemists are generally weak. It is certainly is not what is taught today. Part of me is now discovering how much junk was fed to us. This is why you will see me questioning the very basics once again-so that I can correct myself. $\endgroup$
    – AChem
    Sep 22, 2020 at 1:40

The problem with crystallization of a supersaturated solution is that all molecules are equal. No particular molecule is "entitled" to be the first to start the first crystal. An external event is needed to "nominate" the place of the first crystallization. This external event can be an impurity, a sudden shock, an irregularity at the surface of the container. But once the first face of the first microcrystal is created, all other molecules "know" where and how to develop the crystallization process.

  • $\begingroup$ What physically happens to the molecule that allows it to become the first mircocrystal? $\endgroup$ Sep 20, 2020 at 21:56
  • 1
    $\begingroup$ In the aqueous solution, the ions are surrounded by adsorbed water molecules. These molecules must be removed for making the first arrangement of ions like in the crystal. A sudden shock may be enough to expel for a while these adsorbed water molecules around a given ion. Once the first line of ions is done, the rest is following. $\endgroup$
    – Maurice
    Sep 21, 2020 at 10:56

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