Eddygp
National Hazard
Posts: 858
Registered: 31-3-2012
Location: University of York, UK
Member Is Offline
Mood: Organometallic
|
|
Unknown reaction
I had thought about reacting sodium bicarbonate and magnesium sulfate in aqueous solution. I had expected it to form sodium sulfate and magnesium
bicarbonate or some sort of carbonate and bisulfate. However, I was surprised when I saw some CO2 bubbles forming in the solution. What did I make?
Thanks in advance.
there may be bugs in gfind
[ˌɛdidʒiˈpiː] IPA pronunciation for my Username |
|
|
LanthanumK
Hazard to Others
Posts: 298
Registered: 20-5-2011
Location: New Jersey
Member Is Offline
Mood: No Mood
|
|
Magnesium bicarbonate forms but is unstable, easily decomposing to magnesium carbonate and carbon dioxide.
hibernating...
|
|
|
plante1999
International Hazard
Posts: 1936
Registered: 27-12-2010
Member Is Offline
Mood: Mad as a hatter
|
|
You forgot the water!
2NaHCO3 = MgSO4 -) MgCO3 + Na2SO4 + CO2 + H20
I never asked for this.
|
|
|
Eddygp
National Hazard
Posts: 858
Registered: 31-3-2012
Location: University of York, UK
Member Is Offline
Mood: Organometallic
|
|
But... is Na2SO4 hence more stable than NaHCO3?
there may be bugs in gfind
[ˌɛdidʒiˈpiː] IPA pronunciation for my Username |
|
|
cyanureeves
National Hazard
Posts: 744
Registered: 29-8-2010
Location: Mars
Member Is Offline
Mood: No Mood
|
|
but will sodium sulfate really be formed as well?the same hard rock stuff left over after nitric acid synthesis?just like that?
|
|
|
sargent1015
Hazard to Others
Posts: 315
Registered: 30-4-2012
Location: WI
Member Is Offline
Mood: Relaxed
|
|
That is the correct reaction. I use this all the time to make magnesium carbonate as a precursor to all the other magnesium salts you could ever want.
Realize that the sodium sulfate will be in the aqueous solution since sodium salts are incredibly soluble. You could filter off the Magnesium carb and
then crystallize the mother liquor left over for sodium sulfate if you wanted.
|
|
|
Eddygp
National Hazard
Posts: 858
Registered: 31-3-2012
Location: University of York, UK
Member Is Offline
Mood: Organometallic
|
|
Umm... I do not have any precipitate  It's all dissolved...
EDIT: Sorry, it may be MgCO3·5H2O, which is soluble
[Edited on 19-5-2012 by Eddygp]
there may be bugs in gfind
[ˌɛdidʒiˈpiː] IPA pronunciation for my Username |
|
|
Nicodem
Super Moderator
Posts: 4230
Registered: 28-12-2004
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by cyanureeves  | | but will sodium sulfate really be formed as well?the same hard rock stuff left over after nitric acid synthesis?just like that?
|
No, it is impossible for the sodium sulfate to form in such a reaction. This is aqueous precipitation chemistry and thus the reaction equation that
plante1999 wrote is correspondingly wrong.
The reason why no sodium sulfate can form is because this salt dissociates in water. Furthermore, it can not be directly crystallized from aqueous
solutions unless heating to high temperatures to drive off the water from its hydrates. Therefore, the only sodium compound that could form in this
reaction, but only if the solutions are concentrated enough to reach saturation, is the sodium sulfate decahydrate. This only has moderate solubility
at ambient temperatures and can easily form large crystals. Binary sulfates of the two metals also exist, but should not form under proper reaction
stoichiometry.
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
|
|
|
plante1999
International Hazard
Posts: 1936
Registered: 27-12-2010
Member Is Offline
Mood: Mad as a hatter
|
|
| Quote: Originally posted by Nicodem | | Quote: Originally posted by cyanureeves | | but will sodium sulfate really be formed as well?the same hard rock stuff left over after nitric acid synthesis?just like that?
|
No, it is impossible for the sodium sulfate to form in such a reaction. This is aqueous precipitation chemistry and thus the reaction equation that
plante1999 wrote is correspondingly wrong.
The reason why no sodium sulfate can form is because this salt dissociates in water. Furthermore, it can not be directly crystallized from aqueous
solutions unless heating to high temperatures to drive off the water from its hydrates. Therefore, the only sodium compound that could form in this
reaction, but only if the solutions are concentrated enough to reach saturation, is the sodium sulfate decahydrate. This only has moderate solubility
at ambient temperatures and can easily form large crystals. Binary sulfates of the two metals also exist, but should not form under proper reaction
stoichiometry. |
I did not bother to wright the hydratation of each reactant....
When I do a synthesis I calculate it since It could make all the reaction go wrong if stoichiometry is not respected.
I never asked for this.
|
|
|
weiming1998
National Hazard
Posts: 616
Registered: 13-1-2012
Location: Western Australia
Member Is Offline
Mood: Amphoteric
|
|
Since the reaction is in aqueous solution, just combining solutions of NaHCO3 and MgSO4 will result in a mix of Mg2+ ions, HCO3- ions, SO4- ions, and
Na+ ions. I tried to find data on the solubility of Mg(HCO3)2 but I can't (all I got was some scam cancer treatment). Mg(HCO3)2 is unstable, but the
HCO3- ion itself is fairly stable, and won't decompose in aqueous solution at normal pH. So just combining dilute aqueous solutions of MgSO4 and
NaHCO3 will just result in a pool of ions. Combining concentrated/saturated solutions, however, might cause some Mg(HCO3)2 to fall out of solution if
its solubility is lower than sodium bicarbonate's, these might decompose to MgCO3, resulting in the bubbles you see. If Mg(HCO3)2's solubility is
higher than NaHCO3, however, limited or no reactions occur unless you pour water directly on to solid MgSO4 and NaHCO3.
Magnesium carbonate pentahydrate, although slightly soluble, is still relatively insoluble (http://en.wikipedia.org/wiki/Magnesium_carbonate), so precipitation will still happen, maybe just too little to see compared to the CO2 gases with
a much larger volume.
|
|
|
99chemicals
Hazard to Others
Posts: 174
Registered: 24-3-2012
Location: In the Octet
Member Is Offline
Mood: No Mood
|
|
That is not the correct reaction. The correct equation would be two steps:
(2)NaHCO3 + MgSO4 --> Mg(HCO3)2 + Na2SO4
Then
Mg(HCO3)2 --> MgCO3 + H2O + CO2
Nicodem was talking about the solubility of sodium sulfate in a hydrated form. I thought that when a salt with waters of crystalization goes into
solution it gives the water to the solution and becomes ions. Is he confused or am I?
|
|
|
sargent1015
Hazard to Others
Posts: 315
Registered: 30-4-2012
Location: WI
Member Is Offline
Mood: Relaxed
|
|
| Quote: Originally posted by 99chemicals |
(2)NaCO3 + MgSO4 --> Mg(HCO3)2 + Na2SO4
Then
Mg(HCO3)2 --> MgCO3 + H2O + CO2
Nicodem was talking about the solubility of sodium sulfate in a hydrated form. I thought that when a salt with waters of crystalization goes into
solution it gives the water to the solution and becomes ions. Is he confused or am I? |
My bad, I was thinking about the slightly different reaction:
MgSO4 + Na2CO3 --> MgCO3 + Na2SO4
(which is however not quite right, see link below)
The sodium carbonate can be generated in situ by heating up the sodium bicarb reaction with magnesium sulphate. THIS is what I did to make
my magnesium carbonate.
Alternatively, you could just make sodium carbonate ahead of time and use it.
Check out this link for the reaction/procedure
http://www.nuffieldfoundation.org/practical-chemistry/making...
|
|
|
Eddygp
National Hazard
Posts: 858
Registered: 31-3-2012
Location: University of York, UK
Member Is Offline
Mood: Organometallic
|
|
I will probably crystallize it first and then... try to see what it is.
there may be bugs in gfind
[ˌɛdidʒiˈpiː] IPA pronunciation for my Username |
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
