It can sometimes be difficult to visualise chemistry, with seemingly infinitely small atoms going on about their business at a speed that can be too fast or, more often, too slow to fully appreciate. The Hot Ice demonstration takes on this challenge, taking advantage of the supersaturation that can occur in a sodium ethanoate solution to show, without a chemical reaction even occurring, its sudden crystallisation that propagates as a wave through the solution. It does this while releasing substantial heat, giving the otherwise bizarre effect of heating a liquid which then subsequently solidifies. A very easy and satisfying demo!
Many organic molecules can form supersaturated solutions, ‘supersaturated’ meaning the state of a solution that contains more of the dissolved material than could be dissolved by the solvent under normal circumstances (source: Wikipedia). Sodium ethanoate, also known as sodium acetate, can be easily manipulated to achieve this – a hot saturated solution should be just left undisturbed and allowed to cool down.
As the solution cools down, the water can hold less and less sodium ethanoate dissolved, but, as long as there is a lack of nucleation points – cracks and spaces that would allow crystals to form – all of the compound will stay dissolved. If enough nucleation points are present, such as some dust falls in or a small crystal of sodium ethanoate is dropped in, the dissolved sodium ethanoate will form crystals on them, causing a run-away crystallisation of the whole solution, turning from a thin liquid to a hard mass of solid crystals.
This would be classed as a physical reaction, rather than a chemical one, as no chemical reaction is actually happening – just the switching of physical states as the solid is formed. This is where the heat is generated. Most liquids, when they freeze to become solids, release some energy as bonds form between the chemical to make the solid. Yes, even water does this! When water freezes it releases a small amount of energy known as the enthalpy of fusion due to bonds forming between the water molecules.
This is similar to what is happening when the sodium ethanoate crystallises, where the heat being released is due to bonds forming between the sodium ethanoate molecules as the crystals form, although this energy is more substantial.
|Sodium ethanoate||Low hazard|
- Gloves and goggles should always be worn when handling chemicals to protect your sensitive eyes and skin.
- This is a very safe demonstration, your only worries are really just dealing with high temperature water as you slowly boil the sodium ethanoate solution down, making sure the sodium ethanoate does not begin to decompose at too high temperatures, or burn if the beaker is accidentally boiled dry.
As mentioned, this is a very simple demonstration to set up. A 250ml glass beaker was placed on a hotplate stirrer, equipped with a stir bar.
I added an arbitrary amount of sodium ethanoate trihydrate that I had made a while ago to the beaker, maybe ~10g. The amount needed really does not matter; even this small amount with be enough to give a good effect. Do remember to leave a little sodium ethanoate left though to start the reaction later on.
Enough distilled water should then be added to dissolve the amount of sodium ethanoate at 90ºC plus a bit extra to account for the loss of any water by evaporation while getting up to this temperature. For 10g of sodium ethanoate 15ml of distilled water should be sufficient.
The heating of the hotplate was turned on, as well as medium-strength stirring to reduce the chance of any decomposition of the sodium ethanoate. The temperature was monitored with a thermometer to make sure it was kept under, and then eventually at, 90ºC.
Once achieving 90ºC, the solution was left at this temperature to let the excess water boil off until a thick white solid (the sodium ethanoate) started precipitating around the edge of the solution touching the beaker. A little distilled water, a few drops at a time, are then added to attempt to redissolve the white precipitate.
Once it is fully redissolved, the beaker should be removed from heat and placed to the side to cool slowly. The thermometer and stir bar can be quickly removed, but I decided to leave the stir bar in, seeing if it caused the crystals to form in a circle around the stir bar as if running around a race track. A small piece of cling-film is recommended to put on top of the beaker to keep out any foreign matter, such as dust, that may start the demonstration prematurely.
Once the beaker has cooled down to room temperature, the solution is ready for the demonstration.
The cling-film was removed and a small sodium ethanoate crystal was dropped into the solution. I consequently filmed the reaction, afterward taking a screenshot every exact second – this way you can get an appreciation for the speed of the demonstration.
I will leave you with these pictures now with the odd few words. I hope you do appreciate how quickly these crystals form, especially compared to the copper(II) sulfate(VI) crystal I grew for my blog banner, taking 8 months or so to grow the same distance these sodium ethanoate crystals did in 10 seconds. Now that is impressive!