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Texium
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Quote: Originally posted by Brain&Force  | Yellow tint? How'd you source your mineral oil - could something be coordinating to it?
I've had no problems storing ferrous sulfate in a sealed vial. Purging the vial was found to be unnecessary. |
I'm starting to realize that now, as the stuff that I had sitting out dry open to the air is completely clean. I'm making a fresh
batch. My mineral oil came from Walgreen's.
Also, I noticed that there seem to be some side reactions involved with making it that don't seem to be pointed out usually. This time I tried using
stoichiometric amounts of each reactant, and there was a lot of iron left over, as well as the presence of quite a bit of sulfur dioxide. I think that
one or both of the following equations might have been occurring in addition to the expected one:
2Fe + 3H2SO4 --> Fe2O3 + 3SO2 + 3H2O
Fe + H2SO4 --> FeO + SO2 + H2O
Please correct me if I am blatantly mistaken. 
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blogfast25
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| Quote: Originally posted by zts16 |
2Fe + 3H2SO4 --> Fe2O3 + 3SO2 + 3H2O
Fe + H2SO4 --> FeO + SO2 + H2O
Please correct me if I am blatantly mistaken. |
Unfortunately you are.
Iron reacts with sulphuric acid (at least 'dilute') to form ferrous sulphate (FeSO<sub>4</sub> and hydrogen.
Fe is incapable of reducing sulphate ions (oxidation number +6) to sulphur dioxide (oxidation number +4).
Once crystallised to the heptahydrate, the ferrous sulphate does tend to oxidise by air oxygen to ferric sulphate
(Fe<sub>2</sub>(SO<sub>4</sub><sub>3</sub> as well as ferric hydroxy sulphates and even ferric hydroxide, but those reactions are
a rather slow process.
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Amos
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zts16, just make sure the environment the crystals are in is at a low pH, preferably maintained by sulfuric acid. I've had crystals left in sulfuric
acid solution for over a month, exposed to the air, and not a hint of oxidation has been detected.
[Edited on 9-23-2014 by No Tears Only Dreams Now]
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blogfast25
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| Quote: Originally posted by No Tears Only Dreams Now | zts16, just make sure the environment the crystals are in is at a low pH, preferably maintained by sulfuric acid. I've had crystals left in sulfuric
acid solution for over a month, exposed to the air, and not a hint of oxidation has been detected.
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Give it time: eventually you'll see the first signs of oxidation. But it your conditions it is slow, for sure.
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Texium
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| Quote: Originally posted by blogfast25 | | Quote: Originally posted by zts16 |
2Fe + 3H2SO4 --> Fe2O3 + 3SO2 + 3H2O
Fe + H2SO4 --> FeO + SO2 + H2O
Please correct me if I am blatantly mistaken. |
Unfortunately you are.
Iron reacts with sulphuric acid (at least 'dilute') to form ferrous sulphate (FeSO<sub>4</sub> and hydrogen.
Fe is incapable of reducing sulphate ions (oxidation number +6) to sulphur dioxide (oxidation number +4).
Once crystallised to the heptahydrate, the ferrous sulphate does tend to oxidise by air oxygen to ferric sulphate
(Fe<sub>2</sub>(SO<sub>4</sub><sub>3</sub> as well as ferric hydroxy sulphates and even ferric hydroxide, but those reactions are
a rather slow process. |
So what then is causing the smell of sulfur dioxide? It's quite strong: far more
pungent than the faint fumes normally present over concentrated sulfuric acid. I kept the reaction quite cool to avoid boiling away the acid. The acid
that I have is lab grade, so there shouldn't be anything else in there other than the acid, the iron, and whatever else was in the nail which
shouldn't be very much.
I'm not saying this in an argumentative way, but rather because it still puzzles me.
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Brain&Force
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There are often sulfur impurities in iron, and adding acid in an oxidizing environment will release not sulfur dioxide, but hydrogen sulfide.
At the end of the day, simulating atoms doesn't beat working with the real things...
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Texium
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I've smelled both of those gases before, and this is definitely mostly sulfur dioxide. There might be some hydrogen sulfide, but only a trace amount.
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Amos
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If you're trying to compare sulfuric acid and sulfur dioxide smells, I'm not sure you should do so. They both smell very different, and I believe that
it is sulfur TRIoxide coming off of the acid, but you can correct me if I'm wrong. Iron sulfide impurities are fairly easily oxidized to ferrous
sulfate at higher temperatures.
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blogfast25
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| Quote: Originally posted by zts16 | [ It's quite strong: far more pungent than the faint fumes normally present over concentrated sulfuric acid. I kept the reaction quite cool to avoid
boiling away the acid. The acid that I have is lab grade, so there shouldn't be anything else in there other than the acid, the iron, and whatever
else was in the nail which shouldn't be very much.
I'm not saying this in an argumentative way, but rather because it still puzzles me. |
How strong is your concentrated sulphuric acid? 95? 98? Higher?
And why use conc. H2SO4 to dissolve iron, when dilute works perfectly fine?
Dissolving most metals in acid does produce strange, sulphurous smells. A nail is of course not pure iron either...
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Texium
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I didn't use concentrated sulfuric acid. I have 93%, and I diluted a stoichiometric amount of it. I'm not sure exactly what concentration I diluted it
to.
After it didn't fully react, I doubled the amount of acid and the reaction proceeded, but the nail was still not fully dissolved.
I dissolved another nail in HCl, and it reacted completely and there was only a faint odor of hydrogen sulfide.
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blogfast25
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| Quote: Originally posted by zts16 | I didn't use concentrated sulfuric acid. I have 93%, and I diluted a stoichiometric amount of it. I'm not sure exactly what concentration I diluted it
to.
After it didn't fully react, I doubled the amount of acid and the reaction proceeded, but the nail was still not fully dissolved.
I dissolved another nail in HCl, and it reacted completely and there was only a faint odor of hydrogen sulfide. |
As I wrote above, with diluted sulphuric acid and iron you can't get SO2: there's simply nothing there to reduce the sulphate ions to sulphurous
oxide.
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Little_Ghost_again
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Fabulous crystals! Another experiment on my to do in the future list! Interesting reading as well. I have great respect for the shear amount of
knowledge floating around on here.
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Texium
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| Quote: Originally posted by blogfast25 | | Quote: Originally posted by zts16 | I didn't use concentrated sulfuric acid. I have 93%, and I diluted a stoichiometric amount of it. I'm not sure exactly what concentration I diluted it
to.
After it didn't fully react, I doubled the amount of acid and the reaction proceeded, but the nail was still not fully dissolved.
I dissolved another nail in HCl, and it reacted completely and there was only a faint odor of hydrogen sulfide. |
As I wrote above, with diluted sulphuric acid and iron you can't get SO2: there's simply nothing there to reduce the sulphate ions to sulphurous
oxide. |
Very well then. I will just assume that the nail is too impure for me to determine what is actually
happening in this reaction and leave it at that.
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aga
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Nope.
Get some totally Pure Iron and do comparative tests.
Scientifically.
Then post the results so we can All learn something new.
[Edited on 24-9-2014 by aga]
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Texium
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| Quote: Originally posted by aga | Nope.
Get some totally Pure Iron and do comparative tests.
Scientifically.
Then post the results so we can All learn something new.
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I plan on doing that to make sure that my sulfuric acid is good. It should be, but you never
really know.
I don't currently have any pure iron unfortunately (hence my use of nails in these reactions)
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Bert
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Threads Merged
8-5-2018 at 07:16 |
Bert
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Thread Split
8-5-2018 at 07:20 |
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