August 08, 2022, 04:27:25 AM
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Topic: Calculation of substance amount of Gases forming out of a powder mixture  (Read 889 times)

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Offline Lastipasti

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Hello,

I'm currently conducting experiments where I place sealed silica ampules containing atmospheric air and different powder mixtures into a vertical tube furnace at 1000°C. Only the bottom parts of the silica tubes are exposed to the 1000°C, while the upper parts of the tubes are as cool as ~100°C during my syntheses. The goal of my experiments is to degas these powder mixtures, which then percolate upwards as gases in the ampules and then recrystallize on the silica tube's surface again as they cool down. In some experiments I'm facing problems however, that the gas pressure is too big for these ampules and that they're destroyed upon quenching. For that reason, I want to estimate the maximum gas volumina and gas pressures that could possibly be reached in those ampules. For my approach, I have defined the major expected gas species in the ampules and know the mass of my starting powder mixture, which are presented below:

CuSO4*5H2O = 0,3895 g
CuO = 0,0414 g
NaCl = 0,0304 g
KCl = 0,0388 g

Expected gases: H2O(g), H2S(g), O2(g), SO2(g), SO3(g), HCl(g), NaCl(g), KCl(g) and CuCl2(g)

I know that you can calculate the gas pressure of an individual gas phase based on the ideal gas law, but for that I need the substance amount of the respective gas phase. However, I don't know how much of these gases form in the given volume and what their respective volume percentages are, so I assume that 1. All of my starting material degases and 2. That I can calculate for every gas species individually its maximum substance amount (meaning: If I want to consider only the substance amount of O2, I use up all of the oxygen in my CuSO4*5H2O and all of the oxygen in my CuO). Now what I do not know however, is how to calculate the total substance amount based on the mass of my starting powders. I know that you can calculate it normally if you divide the mass by its molecular weight, however, I'm not quite sure how to consider the oxygen in CuSO4*5H2O on its own. My calculation for the O in the copper sulfate for example looks as followed:

1. CuSO4*5H2O = 249,68 g/mol and O = 16 g/mol
2. CuSO4*5H2O = 9 O or 4,5 O2
3. 4,5 * 32(molar mass O2)= 144 g/mol
4. (144 g/mol / 249,68 g/mol)*100 = 57,67%
5. 0,3895 g * 0,5767=0,224625 g for O2
6. 0,224625 g / 32 g/mol = 0,00701953125 mol

Is this calculation correct? And how do I combine the substance amount of O2 from CuSO4*5H2O and the O2 from CuO? Do I just add them up?

I hope my question has been understood. I'm coming from a geoscientific background so I have unfortunately little knowledge on some basic chemistry. So any help would be appreciated!

Best regards,
Lastipasti

« Last Edit: June 13, 2022, 12:30:50 PM by Lastipasti »

Offline Corribus

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There's a lot to unpack in there and it's not entirely clear to me what you're doing. My first question would by why you think that, e.g., NaCl becomes a gas at 1000 degrees C when it has a boiling point of ~1500 C. Also why you think gas pressure is causing your silica tubes to break when they are "quenched". How do you know it's simply not mechanical stresses caused by large temperature differentials?

I think some of your phrasing is not used correctly as well, which makes it hard to tell what your goals are. For instance., "degas these powder mixtures". Do you mean somehow to purify them by sublimation?
What men are poets who can speak of Jupiter if he were like a man, but if he is an immense spinning sphere of methane and ammonia must be silent?  - Richard P. Feynman

Offline Lastipasti

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Hello Corribus,

thank you for your answer. The ultimate goal of my experiments is to partly recreate the geochemical environments/conditions found in natural fumaroles, as I have a geoscientific background. In fumaroles, you have very complex gas mixtures which can be 1000°C degrees and hotter at its source, which is usually a magma. As the gases are released from the magma, they percolate upwards in the crust to interact with the surrounding rocks and the atmosphere. The further they travel away from the magma source, the more the gas mixtures cool down causing some of the gas complexes to deposit as crystalline phases at pretty much the whole temperature range between 1000°C and 100°C.

Now it is correct, that NaCl on its own should not be subliming at these temperature conditions, however, with the other present phases, I assume it did react to complexes that are much less temperature stable. This assumption is based also on various observations of the silica tubes, where I found no powders or melt phases remaining at the bottom of the tubes. In natural fumarole environments, we see a similar behaviour, since minerals like halite (NaCl) or sylvine (KCl) also deposit from the gases. Due to the complexity of the gas mixtures however, it is difficult to tell which complexes are directly involved in these deposition processes though.

I believe the gas pressure to be causing this, since other experiments were not affected at all by the quenching of the silica tubes (meaning, I put the hot silica tube directly into cold water and they did not break). Only silica tubes of syntheses, where a significant amount of H2O remained in the starting material, exploded on me. So I assume its a combination of increased gas pressure and mechanical stress that makes the silica tubes explode. These tubes also have a very low volume of around 5.5 to 6 cm^3.

When I say degassing, which is a commonly used term in the geosciences, I mean that I want to turn most or all of the crystalline starting materials either into a vapor or a gas phase so that they're transported away from the hot zone of the tube. If you want to know what the set-up looks like, it is similar to the one described in this work: https://agupubs.onlinelibrary.wiley.com/doi/full/10.1029/2018JE005911

Compared to the set-up described in this work however, I don't use a melt as my source for volatiles (or gases) but instead I expect my whole source material to sublime. I hope I explained this a bit better now.

Online Borek

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You should be able to do some estimates just by looking at the possible stoichiometry. For example if you have 0.3895 g of copper sulfate you can easily calculate number of moles of water and number of moles of S present. Sum of numbers of moles of SO2, SO3 and H2S must equal the initial number of moles of S - plus, actually the gas composition doesn't matter much, as it is number of moles of gas that matters, and that you know.

I would just write mass balances and play with different compositions by trial and error, looking for the worst case. Still, my bet is it is the water that is the main culprit.
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Offline Lastipasti

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Hello Borek,

thanks for that comment. I was thinking about doing such calculations, however, I wasn't sure if it was the optimal way to approach this problem. And as for your guess, I agree that the water is the main culprit. Problem for me however is, that since I want to simulate natural fumarolic environments in nature where water is always occuring in abundance, I need to find a balance between the maximum mass of powder I can use vs. the maximum amount of gas pressure my silica tubes can handle. But I'll try it the way you suggested and see about the results!


Online Borek

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I wasn't sure if it was the optimal way to approach this problem

In theory it should be possible to calculate the equilibrium composition of the mixture from thermodynamic data (I believe there are programs written for this purpose by your geochemistry colleagues, sorry, don't remember any name). Unfortunately these 1. are not very accurate, 2. describe theoretical equilibrium, not the actual situation influenced by kinetics, 3. because of 1. and 2. won't actually give substantially better answers than those from just mass balance evaluations.
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