sbreheny - 30-1-2014 at 00:15
Hi all,
I have two questions regarding hydrates/water of crystallization that I have not been able to find an answer to via "The Google":
1) I am trying to convert Cobalt Chloride Hexahydrate to the anhydrous form. I believe that this can be done by heating but I wanted to try doing it
by vacuum desiccation. After a few hours under vacuum with some silica gel, it converted from the hexahydrate to what seems to be the dihydrate (both
by appearance and by the amount of mass change). I do not seem to be able to convert it further to the anhydride as it has been about 24 more hours
under vacuum and there has been no further mass change or appearance change. Can anyone point me to some info source on the limitations of different
drying methods for various different hydrates which might enlighten me as to why I have hit this limit?
2) Various sources online list the melting point of hydrates as much lower than their anhydrous form (for example, Calcium Chloride dihydrate versus
anhydrous). When I tried to test this by heating, I simply drove out the water without any melting going on. How would one go about observing the
dihydrate melting when heating it drives off the water first?! Do you have to also control the humidity as you heat to preserve it in the dihydrate
form? If so, it seems like such melting point data is rather useless in practice as you are unlikely to have a constant humidity as the temperature
changes. Is the melted dihydrate actually more like simply an aqueous solution of CaCl2 or does it actually look like molten CaCl2?
Thanks!
Sean
woelen - 30-1-2014 at 00:25
Cobalt chloride anhydrous can be converted to the anhydrous form by careful heating:
http://woelen.homescience.net/science/chem/exps/raw_material...
The "melting of hydrates" is not really melting, it is formation of a solution of the ionic compound in the water of crystallization. At increasing
temperature, for many ions, the level of hydration decreases and hence, molecules of water become available for dissolving the less hydrated ions.
Many hydrates "melt" well below 100 C. Anhydrous salts, if they do not decompose, do real melting, but this usually is at very high temperatures (at
least several hundreds of degrees Centigrade, some salts even need temperatures near 1000 C).
Oscilllator - 30-1-2014 at 00:25
with regards to 1) I understand that the more water ligands you pull off a metal ion, the more strongly the remaining ligands are held by the metal
ion. Therefore, the vacuum may only be strong enough to remove 4 water ligands and you may either have to a) increase the vacuum or, more likely, b)
apply heat to the cobalt chloride. If the Cobalt chloride is under vacuum than it may not be necessary to apply a lot of heat - a water bath may be
able to do the trick.
blogfast25 - 30-1-2014 at 03:50
Salt/hydrates/watery solutions v temperature two phase diagrams help understand which hydrate can exist in equilibrium with at sat. solution at which
temperature:
https://www.google.co.uk/search?sourceid=navclient&ie=UT...
The second link to a *.pdf on soda shows the Na2CO3/water two phase diagram and explains it all fairly well.
[Edited on 30-1-2014 by blogfast25]
sbreheny - 1-2-2014 at 23:20
Thanks to all who responded. I finally was able to get anhydrous CoCl2 from the Hexahydrate. I began by heating it in a water bath to 90 C under a
partial vacuum (25 inHg was the best I could do). Once again, this converted it to the dihydrate, although much more rapidly this time than it was
with vacuum/desiccant and no heat. However, 90 C was not hot enough to go the final step to the anhydrous form. I removed it from the water bath and
stopped the vacuum and just heated the flask on a hotplate - to around 200C. I did notice that briefly it got too hot and began to evolve acid fumes
so I quickly backed off the heat. Anyway, I now have an azure powder and the ratio of the final mass to the original mass was 0.5434. Theoretical is
0.5457! I'm pretty happy!