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Author: Subject: The Short Questions Thread (4)
FableP
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[*] posted on 12-11-2023 at 21:53


High Powered Electric Propulsion - Electric solid propellant

http://eplab.ae.illinois.edu/Publications/IEPC-2019-421.pdf
and
https://scholarsmine.mst.edu/cgi/viewcontent.cgi?article=383...

Table 1: Chemical composition of the High Performance Electric Propellant (HIPEP).

Hydroxyl Ammonium Nitrate (HAN) (NH3OH)+ NO3 - 75%
Polyvinyl Alcohol (PVA) CH2CH(OH) - 20%
Ammonium Nitrate (AN) NH4NO3 - 5%

Thanks for pointing out the correct formula for barium nitrate.
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j_sum1
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[*] posted on 12-11-2023 at 22:28


Mmm.
I am always a bit concerned when someone who can't balance an ionic formula starts playing around with significant quantities of energetics.

Be aware also that barium compounds are toxic to humans and the environment. Make sure you manage the risks and have a dispisal plan.
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[*] posted on 13-11-2023 at 06:11
Why does the nitrate ion exist?


The obvious answer is, because it does.
But the charge distributions make no sense to me
(I get the idea of resonance but don't understand why the ion is stable)
(is it the surrounding water molecules that stabilise it?)
anyone have an answer that I might understand?




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clearly_not_atara
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[*] posted on 13-11-2023 at 08:43


Symmetry!

Nitrate has three resonances:

(ON+(O-)O- + O-N+(O)O- + O-N+(O-)O) / sqrt(3)

According to the uncertainty principle, the more space an electron orbital occupies, the smaller the associated kinetic energy it requires to be admissible. It's got nothing to do with water molecules or other coordinations; the stability of alkali metal nitrates increases with ion size:

CsNO3 > RbNO3 > KNO3 > NaNO3 > LiNO3

with CsNO3 stable to around 700 C and LiNO3 decomposing below 400 C. Of course, this is precisely opposite to the coordination strength of the ions. And indeed the hydrogen compound is already slowly decomposing at room temperature.




[Edited on 04-20-1969 by clearly_not_atara]
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[*] posted on 6-12-2023 at 14:45
Iron (II) sulfate


I bought some sulfate of iron fertilizer from a garden centre wit the aim of purifying it to some FeSO4 crystals. When dissolved in water there was allot of light brown material which passed through my filter so I left it to settle for 24 hours.

I managed to pour off a nice green coloured solution and discarded the brown sludge.

I then tried to boil some of this down with the idea to drop out some crystals of iron sulfate.

What I notice is that as it is heated and concentrated I get a brown residue on the beaker and my solution become more cloudy. Is Iron II Sulfate difficult to purify? I see it might produce brown oxides in the presence of water.

Would I be better pouring it all into a big bucket and simply allowing it to evaporate over a few weeks to grow crystals?
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[*] posted on 6-12-2023 at 14:52


The brown crud is from oxidation of iron(II) to iron(III), which happily precipitates as various forms of rust. You might try adding excess sulphuric acid to prevent this.



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[*] posted on 6-12-2023 at 15:06


Quote: Originally posted by DraconicAcid  
The brown crud is from oxidation of iron(II) to iron(III), which happily precipitates as various forms of rust. You might try adding excess sulphuric acid to prevent this.


You are of course right, I just added some sulphuric acid to my hot boiling solution and it has gone from brown to green! I am learning!
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6-12-2023 at 15:15
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[*] posted on 6-12-2023 at 17:00


PS- if you're looking to get crystals, rather than just purifying the compound, I believe double salts such as (NH4)2Fe(SO4)2 give better crystals than the plain sulphate.



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Sir_Gawain
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[*] posted on 14-12-2023 at 07:50
Thorium nitrate


Will ThO2 dissolve in nitric acid?

[Edited on 12-14-2023 by Sir_Gawain]




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[*] posted on 15-12-2023 at 00:12


I don't think so, or only with great difficulty. ThO2, like many other oxides, is quite inert if it is calcined.
When it is freshly prepared from aqueous solution, then you have some hydrous form ThO2.xH2O, which easily dissolves in acids. These hydrous oxides (or mixed oxides/hydroxides??) are quite different from the anhydrous calcined oxides.




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Sir_Gawain
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[*] posted on 26-12-2023 at 15:51


Could molecular sieves be used to dry nitric or sulfuric acids? I can’t think of a reason why it wouldn’t work for nitric, but I’m probably missing something.

[Edited on 12-26-2023 by Sir_Gawain]




“Alchemy is trying to turn things yellow; chemistry is trying to avoid things turning yellow.” -Tom deP.
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[*] posted on 26-12-2023 at 16:02


Never mind. I just tried it with some azeotropic nitric acid and it completely destroyed the sieves.



“Alchemy is trying to turn things yellow; chemistry is trying to avoid things turning yellow.” -Tom deP.
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