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Context

I am trying to synthesize and isolate boric acid from borax. I know that boric acid is less soluble in water than the NaCl that will be in the solution so it will precipitate out more readily, but I want to improve the yield by separating the additional dissolved boric acid from the solution. To do this I planned on completely drying the remaining solution and using (near) anhydrous isopropyl alcohol to dissolve the boric acid while leaving the NaCl. Because I didn't have any high concentration isopropyl alcohol around, I decided to use my anhydrous magnesium sulfate to concentrate it.

Afterwards I was going to dry it again which I did by first by lighting the alcohol on fire on a tray (which got it mostly dry) and then putting it in my oven (on broil for <10 min; I typically do it for longer at a lower temperature when going from the heptahydrate to the anhydrous (because I need to go slowly when there is that much water in it to not make a mess and to avoid caking). That being said, I assume my oven does not reach temperatures anywhere near as high as 1,124°C (the decomposition temperature for anhydrous MgSO4) because the heating element is still red (not yellow like one would expect a metal to appear at around 1,000°C) aluminum foil doesn't melt in it, and information on the internet which corroborates this: https://genuineideas.com/ArticlesIndex/cookieBakingSimulation.html).

Despite this, when I took it out I could smell sulfur trioxide, and the appearance of the powder had changed from white to grey. I dissolved it into water (was exothermic as expected) and the solution was definitely acidic with suspended, now dark grey particles which I suspect to be magnesium metal particles covered in a layer of MgOH because of their grey color and apparent low density (as they remain suspended for a long time).

Question

All of this leaves me fairly confident that after the typical dehydration decomposition reaction chain, my anhydrous MgSO4 did decompose somehow, but leaves me wondering how it happened in my standard oven.

Does anyone have a better idea what is going on here? Thanks in advance for your answers.

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  • $\begingroup$ Decomposition of magnesium sulfate will give magnesium oxide, which forms a sparingly soluble hydroxide but is measurably basic. Not acidic. $\endgroup$ Commented Nov 16, 2023 at 2:00

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OP's question:

[...] fairly confident that my $\ce{MgSO4}$ did decompose somehow, but leaves me wondering how it happened in my standard oven.

I'm sure OP's standard oven can heat up to $\pu{500 ^\circ C}$ even though OP didn't say how high it can go. However, according to Ref.1:

Magnesium sulfate hydrate will dehydrate below $\pu{500 ^\circ C}$ to form anhydrous magnesium sulfate. Different temperatures for dehydration are reported in the literature: $330$-$\pu{365 ^\circ C}$ (Ref.2), $400$-$\pu{500 ^\circ C}$ (Ref.3) and at least at $\pu{450 ^\circ C}$ (Ref.2). Magnesium sulfate will decompose into magnesium oxide and sulfur trioxide when the temperature is $\pu{700 ^\circ C}$ or higher, depending on the atmosphere under which the decomposition is carried out. Sulfur trioxide is not stable at temperatures in excess of $\pu{700 ^\circ C}$ and spontaneously forms sulfur dioxide and oxygen. $$\ce{MgSO4 -> MgO + SO3(g)}$$ $$\ce{SO3(g) -> SO2(g) + 1/2O2(g)}$$

Also, keep in mind that the presence of some impurities such ad elemental $\ce{C}$ and gaseous $\ce{CO}$, which can be act as reducing agents will accelerate the decomposition process at lower temperatures (Ref.1):

$$\ce{MgSO4 + 1/2C -> MgO + SO2(g) + 1/2CO2}$$ $$\ce{MgSO4 + CO ->[\pu{600 ^\circ C}] MgO + SO2(g) + CO2}$$

References:

  1. Madeleine N. Scheidema and Pekka Taskinen, "Decomposition Thermodynamics of Magnesium Sulfate," Ind. Eng. Chem. Res. 2011, 50(16), 9550–9556 (DOI: https://doi.org/10.1021/ie102554f).
  2. H. -H. Emons, G. Ziegenbalg, R. Naumann, and F. Paulik, "Thermal decomposition of the magnesium sulphate hydrates under quasi-isothermal and quasi-isobaric conditions," Journal of thermal analysis 1990, 36(16), 1265–1279 (DOI: https://doi.org/10.1007/BF01914050).
  3. M. Seeger, W. Otto, W. Flick, F. Bickelhaupt, and O. S.Akkerman, "Magnesium Compounds," In Ullmann’s Encyclopedia of Industrial Chemistry, Electronic Release; Wiley-VCH: Weinheim, Germany, 2005 (ISBN 13: 9783527310975).
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  • $\begingroup$ Maybe I wasn't clear enough in that end question, but as suggested by the title and my context section, I am referring to decomposition of anhydrous magnesium sulfate. For the context section, tl;dr of what you seemed to miss is this: I describe getting a grey powder, which when I added water and did various basic tests, seems to be made of various (non-magnesium sulfate) magnesium compounds compounds and contained small amounts of sulfur trioxide (which formed a low concentration sulferic acid solution when the water was added). $\endgroup$ Commented May 20, 2023 at 8:18
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    $\begingroup$ And thank you for the section of your post on reducing agents, it certainly is relevant to my question. Do you know any reactions that would be plausible in my situation? I have updated my question to try to make it more clear. $\endgroup$ Commented May 20, 2023 at 8:46

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