chemkid
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Sodium Thiosulfate Melting
Long time no posts...Ive still been doing some chemistry, especially using OTC materials for spinach chromatography.
Anyways,
I wanted to perform an experiment with super cooling and found one using sodium thiosulfate. A melting point of 117 degrees fahrenheit (47.22) seemed
simple enough. So i boiled water to 60 degrees celsius and placed a small test tube with about a gram of sodium thiosulfate in it...no melting after
several minutes. I tried heating on an open flame...decomposition. So i went back and checked the experiment procedure (which warned against heating
over an open flame).
The experiment uses sodium thiosulfate dihydrate, where as i have been using anhydrous sodium thiosulfate. Is this why it refuses to melt? If so, how
can i make it a dihydrate?
Thank you,
chemkid
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Nerro
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Weigh it, do the math and add the correct amount of water.
#261501 +(11351)- [X]
the \"bishop\" came to our church today
he was a fucken impostor
never once moved diagonally
courtesy of bash
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Xenoid
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| Quote: | Originally posted by chemkid
The experiment uses sodium thiosulfate dihydrate, where as i have been using anhydrous sodium thiosulfate. Is this why it refuses to melt? If so, how
can i make it a dihydrate?
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Are you sure it is the dihydrate you want? This experiment normally uses the PENTA-hydrate. Sodium thiosulphate normally forms a pentahydrate. Can't
you just make a hot saturated solution using your anhydrous product and cool it to 0oC, the solubility curve is pretty flat but you should get some of
the pentahydrate crystallising out!
You can then melt (~48 oC) your pentahydrate crystals for the experiment.
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pantone159
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| Quote: | Originally posted by chemkid
i have been using anhydrous sodium thiosulfate. Is this why it refuses to melt? |
Yes - The Na thiosulfate is actually not melting, but rather dissolving in its water of crystallization. If you have no water of crystallization,
this cannot happen. Also, it probably is supposed to be the 5-hydrate, that is the common form.
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chemkid
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Alright. Will try adding water to form the PENTA hydrate, my mistake.
Chemkid
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Xenoid
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For completeness, the thiosulphate supercooling procedure is given below.
A supersaturated solution may be seeded by addition of a small crystal. This will create a solid mass of the chemical.
To demonstrate this:
Half fill a glass with crystals of sodium thiosulphate pentahydrate (Na2S2O3.5H2O);
Heat the glass in a bath of hot water until the crystals "melt" to form a transparent liquid;
Filter out any impurities using a funnel and cotton;
Cover the glass and allow the liquid to cool to room temperature;
Shake the glass, or add a seed crystal, whereupon the liquid "freezes" immediately;
Notice that the glass feels warm.
Sodium thiosulphate freezes at a temperature of 48°C. Thus, a solution of this chemical is supercooled when at room temperature (around 20°C). The
shaking of the glass, or the addition of a small crystal, triggers the freezing process. When freezing occurs, latent heat is released and the glass
is therefore warmed.
Those wishing to experiment further with this phenomenon will find it is exhibited by many, if not most, salt hydrates.
I discovered to my surprise that xylitol exhibits supercooling extremely well. Xylitol is a polyol (poly-alcohol) and is used as a sugar substitute in
dietary products. I purchased some for "candy propellent" experiments. It melts below 100 oC. and when cooled to room temperature forms a clear,
glassy, flexible "toffee" like material. I was idly kneading some of this, when it became extremely hot, opaque and crumbly, as it suddenly
recrystallised.
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woelen
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Yes, you MUST use the pentahydrate. However, many sellers of sodium thiosulfate also sell the anhydrous form, I purchased some of this and it is
anhydrous and this nice little experiment does not work for me. If you buy this chemical, then specifically ask which form you get.
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