kyro8008
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Quirky iron sulphate
I made some FeSO4 about a month ago by dissolving iron in sulphuric acid. Anyway, I obtained some beautiful blue/green crystals, I pretty much dried
them and put them in an aluminium storage case. I opened it up today, and they have turned yellow/light brown - horrible colour!
There might have been a small amount of water in there, but very little in any case and I think the container was airtight. Does anyone know what
might have happened here?
Secondly; when I first dissolved the iron, after the reaction had finished there was a very sharp unpleasant irritating slightly metallic pungant
smell, could this have been in part H2S?
Thanks very much!
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12AX7
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I've noticed the smell too, it's probably from minor impurities (typically under 0.05%) of S and P in the alloy (most likely present as iron sulfide
and phosphide, if they aren't actually in solid solution).
I don't like the sound of that "aluminum storage case". Aluminum readily reduces half the periodic table, iron included, and I'm suprised the case
itself didn't burn through or something.
Traces of moisture and oxygen will oxidize FeSO4 to Fe2(SO4)3 and Fe2O3 (or FeOOH or Fe(OH)3). Conversely, lack of moisture will dehydrate the
crystals, turning them to a dry, white (minus impurities) powder.
You'll have to recrystallize to get it nice and cyan again...
Tim
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kyro8008
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Thanks Tim, the "aluminium storage containers" are old film canisters which have come in very useful, I really need to get some more glass containers!
Ill have to think carefully before I put certain chems in them.
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Mr. Wizard
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"Secondly; when I first dissolved the iron, after the reaction had finished there was a very sharp unpleasant irritating slightly metallic pungant
smell, could this have been in part H2S?"
It's very easy to smell iron compounds. They have a metallic smell or they can even smell like blood. I have to put iron sulfide type (Ironite)
compounds on my garden due to the high calcium carbonate levels in the local water, and you can smell and taste the iron after handling it. I always
have to shower afterward. You should really consider using a dust mask or better when handling more than miniscule quantities of any metals salts.
Even solutions can make microscopic droplets when they fizz, bubble ,or boil, carrying the salts into the air. I got a nasty burn in my throat from
handling less than 10 grams of Cupric Acetate, and I was trying not to breath the dust. You don't know what other metals were alloyed with the 'iron'
you dissolved. There could be everything from Copper, Chromium,Manganese, Nickle to Vanadium. Lungs don't regenerate.
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woelen
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Ferrous sulfate is very prone to aerial oxidation and it is really hard to store this for a longer time, without the oxidation to basic ferric
sulfate.
If you want to have a nice iron (II) salt and want to keep it that nice, dissolve some iron in H2SO4 and next add ammonia. The liquid still should
remain fairly acidic. In that situation you have ammonium sulfate and ferrous sulfate in solution. Together these crystallize as a beautiful light
green/blue double salt, Fe(NH4)2(SO4)2.6H2O. This is called Mohr's salt. This salt is stable with respect to aerial oxidation and can be kept for a
long time at high purity.
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If you dissolve iron in sulphuric acid, you also may get some SO2 and H2S and not only H2 while the metal dissolves. This effect is very strong, when
dissolving zinc in sulphuric acid, but I can imagine, that iron also is capable of reducing the sulfate ion somewhat. This is a side-reaction, besides
the main reaction where Fe reduces H(+) ions to H2.
I also agree with the strong smell of iron salts. Just for fun, dip your finger in a solution of an iron salt (doesn't matter which one, e.g. FeSO4,
FeCl2, FeCl3) and rub a little and then smell it. This gives a typical iron smell. To me it is not particularly unpleasant, but it is very typical and
unique to iron. Copper has its own unique smell in this way.
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12AX7
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Speaking of smells (dragging things further off topic  ), nickel and copper
metals have weird reactions with your skin. I always get that weird metallic smell after handing coinage (US nickels, dimes and quarters are clad
with 25% Ni, 75% Cu) or copper (wire, tubing, etc.). Coin smells a little different from pure copper, so I'm betting nickel has its own smell. I've
also noticed smells in use (cutting vegetable or meat) from the steel knife I forged (probably 5160 alloy -- not the usual 400 series stainless
commonly used for knives).
Tim
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hinz
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How can I oxidise Fe(II)sulfate to basic Fe(III)sulfate, something like this Fe3(SO4)2(OH)5 X 2 H2O. I need it for the production of sulfur trioxide.
I put an excess of iron screws in some sulfuric acid and heated it strongly. I waited till no more H2 liberated, heated it stronly and crystallised
the FeSO4 out. Everything worked perfect, but now after a week standing at open air the green FeSO4 won't oxidise to the brown Fe(III)sulfate, the
crystalls are still green. I didn't neutralise the solution, it's still sour (pH arround 1). Does anyone know what inhibit the oxidation of my FeSO4 ?
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12AX7
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Well, low pH will... heh! High oxidation states tend to be unstable at low pH. Fe(III) isn't bad but this rule of thumb may make it difficult to
produce. Plus, in a solution with free acid, the Fe(III) is going to go into solution, where it may not be visible against the majority of Fe(II).
A bunch of ammonium or potassium sulfate might help it, to form ferric ammonium or potassium sulfate (the double salt), but that's probably not what
you want, either.
Why don't you just heat the FeSO4 after dehydrating?
Tim
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hinz
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Thanks, ididn't know that Fe(III) is unstable at low Ph.
""Why don't you just heat the FeSO4 after dehydrating?""
I've tried it, but it liberates a mixture of SO3 and SO2 whereas Fe2(SO4)3 liberates mostly SO3.
Fe2(SO4)3 ==> Fe2O3 + 3 SO3 (at not too high temperatures at which SO3 decomposes to SO2 and O2)
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12AX7
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Ah, ok then. Yeah, it would tend to reduce some of the SO3 as it pyrolyzes.
You could, let's see, got any H2O2? Same as oxygen, but faster... Otherwise, you might get an air bubbler (like for a fish tank).
If you want Fe2(SO4)3, add a mole of sulfuric acid to two moles of FeSO4 before adding the H2O2.
The reaction would be:
2FeSO4 + H2SO4 + H2O2 = Fe2(SO4)3 + 2H2O
Tim
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