annaandherdad
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Reduction by H2 of CuO, Fe3O4, Fe2O3
Reduction of CuO by H2 is a demo given in introductory chem classes, with videos on youtube etc. But there are only occasional references to it on
this forum, with no details (I couldn't find any). I had never see it myself, and was curious to try it. Afterwards I also tried reducing Fe3O4 and
Fe2O3 with hydrogen, with some success. The Fe2O3 produced pyrophoric iron (this was a surprise to me). There are quite a few references on this
forum to pyrophoric iron, but most concentrate on the route through ferrous oxalate. Here are some:
http://www.sciencemadness.org/talk/viewthread.php?tid=8245#p...
http://www.sciencemadness.org/talk/viewthread.php?tid=5639#p...
http://www.sciencemadness.org/talk/viewthread.php?tid=2579#p...
I began by heating a pyrex test tube in a blow torch at the round end, gradually blowing out a hole, which I then reduced to the size of a pinhole and
thickened up by further heating. The pinhole was the exit for the H2 gas. Clamping this test tube horizontally, I first put a small pile of CuO
(obtained by heating CuCO3 in a crucible under a blow torch), and fed H2 gas into the test tube. I used scrap Al + cheap HCl to make the hydrogen,
and got a pretty good flow. After all the air was flushed out of the system, I ignited the H2 emerging from the hole to flare it off, and then heated
the glass under the CuO with a small butane burner. When the reaction started it created a red hot layer that "burned" its way through the pile,
until all CuO was reduced to Cu. This was a spectacular reaction that impressed my daughter and a neighbor. Then I removed the heat and blew out
the H2 flame, allowing the H2 to continue to flow until the tube and contents had cooled, because if air comes in contact with the Cu while hot it
will oxidize and revert to CuO. Then I dumped out the Cu metal and saved it.
Later I tried something similar with Fe3O4. This was magnetite, which I had obtained by running a magnet through sand obtained from Death Valley. I
believe most sand has some magnetite in it. This time the reaction did not create a visible "burn", in fact I'm guessing that it is endothermic. So
I heated the Fe3O4 to red heat with a blow torch under the test tube (the glass must have been getting soft). After cooling, I dumped the contents
out, and could see very little difference between the black results and the original (black) Fe3O4, so I was wondering if it had worked. Dropping a
sample into HCl showed that it had: I got immediate production of H2 gas, while HCl reacts only slowly on the Fe3O4. But soon the H2 gas production
ceased, with plenty of black stuff left behind; was it Fe3O4 that was not converted to Fe? I think so, because the stuff left behind was slowly
attacked by the HCl, producing what looks like FeCl2 and FeCl3.
I thought the imperfect conversion might have to do with the size of the grains of Fe3O4, so I tried this again after grinding the Fe3O4 to a finer
powder. This still gave only partial conversion to Fe, however. Perhaps I would have achieved a better conversion if I had kept the hot iron oxide
longer in the hydrogen?
Finally, I tried the same thing with Fe2O3, a very fine powder ("rouge") that I bought from United Nuclear. In this case the red iron oxide gradually
turned black as I heated it under the hydrogen atmosphere (again to red heat with a blow torch). Again I waited until it was cool, then dumped the
product out onto a piece of paper. Suddenly it began to glow, and ignited the paper! I thought I hadn't waited long enough for it to cool, but later
I realized that I had created pyrophoric Fe.
Any other SF Bay chemists?
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AJKOER
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You may wish to make pyrophoric iron by heating Iron Oxalate. Here is the Youtube demo:
http://www.youtube.com/watch?v=adhE1m2vX38
You can heat the Iron or use Bleach (NaClO) to oxidize it. Then repeat your H2 reduction experiment.
To make Iron Oxalate, per Watt's Oxalic acid will dissolve Iron forming Ferrous Oxalate. Or, you can easily make a ferric salt by slowly dissolving Fe
in a mixture of Bleach (NaClO) and Vinegar. This is basically a dilute solution of HOCl and Sodium acetate which, with time, is capable of forming a
reddish brown Ferric Chloride (and acetate) solution. Only react a dilute solution with Oxalic acid being prepared for a vigorous reaction.
2 FeCl3 + 3 H2C2O4 --> Fe2(C2O4)3 + 6 HCl
If the Iron (III) oxalate does not precipitate, dilute the solution and expose to sunlight.
For other paths to Iron Oxalates, see Watt's "A dictionary of chemistry and the allied branches of other sciences", Volume 4, page 258. Link:
http://books.google.com/books?id=RaA-AAAAYAAJ&pg=PA250&a...
Here is an MSDS for those working with Oxalic acid.
http://www.sciencelab.com/msds.php?msdsId=9926346
[Edited on 24-4-2012 by AJKOER]
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watson.fawkes
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Quote: Originally posted by annaandherdad  | | Afterwards I also tried reducing Fe3O4 and Fe2O3 with hydrogen, with some success. The Fe2O3 produced pyrophoric iron (this was a surprise to me).
[...] Perhaps I would have achieved a better conversion if I had kept the hot iron oxide longer in the hydrogen? |
These two issues are related. What you're seeing is hydrogen diffusion into the solid matrix, where it reacts. Gas-solid diffusion
isn't really rapid, even at elevated temperatures, but it gets there eventually. There's a related reduction of oxides with CO, but it doesn't behave
quite the same way because it can't diffuse nearly as readily. As with any gas-solid diffusion, there's often a threshold temperature above which
diffusion proceeds much more rapidly; it corresponds to the relationship between gas size and lattice spacing.
After the hydrogen diffuses, it forms interstitial H2O. That molecule really doesn't want to be there and its effective pressure inside the
lattice is rather high. This exerts a mechanical strain on the lattice, causing lattice atoms to move. The end result is a lot of surface area formed
by escape channels.
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annaandherdad
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Thanks for the replies, both of you, I appreciate all the information. watson, if what you say is true, I should get a larger conversion of oxide to
iron if I just keep the oxide longer in the gas stream. I mean, it's not just the longer time for the gas to diffuse into the solid, but, if it is as
you say, the surface area increases as the process continues. As for the temperature, I doubt if I can go much higher in my glass test tube, because
the glass will melt. My supply of Fe203 is much, much finer (rouge power) than my sample of magnetite, Fe3O4, and it's possible I had a higher
conversion to Fe metal in that case. But since it burned itself up after I got it out, I didn't have a chance to test that. Of course after it
burned it was back to Fe2O3, but it was a coarser and ruddier product than the original rouge.
I read in Wikipedia that a common method of producing iron these days, industrially, is to use a mixture of CO and H2, produced by reacting CH4 and
O2. This is an alternative to the use of CO alone (produced by passing O2 over coke), the traditional method. And the temperatures are much higher
than I can achieve in my test tube.
The pyrophoric process itself is also interesting. Does any iron surface (no matter how large or small the piece) oxidize on the surface when in
contact with air, at room temperature? I mean, without water, the usual condition for the formation of rust. And if so, then is the pyrophoresis
(if that's the right word) of the small grains simply due to the large surface area, hence the large heat production per unit mass? Or is there some
other effect going on? One would think that in a small sample, even if a lot of heat per mass were being produced, that simple conduction would carry
the heat off efficiently and prevent high temperatures from being reached.
Let me put it this way. If all iron surfaces oxidize in the same way, independent of the size of the sample, then there is a certain thickness that
oxidizes in a certain time. And if the grain size is comparable to that thickness, then in the given time the whole sample would oxidize. Whatever
the details, there are some quantitative questions of this sort.
AJKOER, thanks for the info on the oxalate method. I may try that after I'm done with this approach.
I also read that finely divided Cu is pyrophoric, but I saw no sign of that when I reduced my CuO with H2. Maybe the grains were too large, or maybe
they fused together by the heat of the (obviously exothermic) reaction.
I got into this because I had some copper byproducts of some other things I was doing, and my references say it's bad form to just pour soluble Cu
salts down the drain. So I converted them to CuCO3 and filtered that off, and ended up with a fair amount of CuCO3. It was some trouble to do this,
but then I started looking around for things to do with the CuCO3. One thing I tried was heating CuO and C in a crucible, to try to get elemental
copper. I think I did get some, but the process was messy, with excess carbon around. So the reduction by H2 had the appeal of being cleaner.
I may try some more experiments with this, as long as my hydrogen generator is hooked up. It was fun discovering a phenomenon that I didn't know
about, even if it's well known.
Any other SF Bay chemists?
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blogfast25
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| Quote: Originally posted by annaandherdad |
Later I tried something similar with Fe3O4. This was magnetite, which I had obtained by running a magnet through sand obtained from Death Valley. I
believe most sand has some magnetite in it. This time the reaction did not create a visible "burn", in fact I'm guessing that it is endothermic.
|
Probably. A crude calculation at STP with HoF for Fe3O4 and H2O respectively of - 1121 kJ/mol and - 286 kJ/mol (both NIST values) gives for Fe3O4 + 4
H2 === > 3 Fe + 4 H2O an STP value for HoR of -23 kJ/mol, so that's negligible.
The reaction is driven by the removal of steam in the excess hydrogen stream, similar to what happens in blast furnaces where CO is the main reducing
agent.
Interesting experiments!
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watson.fawkes
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| Quote: Originally posted by annaandherdad | I should get a larger conversion of oxide to iron if I just keep the oxide longer in the gas stream. [...] As for the temperature, I doubt if I can go
much higher in my glass test tube, because the glass will melt.
[...]
Does any iron surface (no matter how large or small the piece) oxidize on the surface when in contact with air, at room temperature? I mean, without
water, the usual condition for the formation of rust. And if so, then is the pyrophoresis (if that's the right word) of the small grains simply due
to the large surface area, hence the large heat production per unit mass? |
If you want to get to higher
temperatures, you can use an iron reaction vessel. You certainly won't attack the vessel, indeed, it will get all shiny inside. And unlike
liquid-phase reductions, you won't have to worry about material melting into the surface. You can use black iron pipe parts. In order to make this
work, you'll need the same kind of fine orifice to make a hydrogen jet. You could use a tiny drill, if you've got a good enough drill press and a
rigid clamping system. There's also a technique (though I've never tried it) which is to embed a copper wire in a lead plug and the dissolve out the
copper with nitric acid.
Many metal surfaces spontaneously oxidize, yes. Iron does so, but then the surface passivates, at least partially. Aluminum passivation is far more
effective, true, but iron does so as well, though to a lesser extend.
It is indeed the surface area to volume ratio that generates pyrophory. Physically, this ratio gives an energy ratio between energy release in
oxidation with the heat capacity of the material. So below some characteristic size, you'll get very rapid temperature rise, which just drives the
reaction faster.
One subtlety is that as the temperature rises, radiative heat transfer starts to dominate. Now if a single isolated metal dust particle is introduced
to air, it will take a smaller size to combust than if you have a quantity of such dust. Why? Because the interior volume of the dust heats up, but
net heat loss only happens at the surface of the dust cloud. As long as there's adequate oxygen, a denser cloud become pyrophoric at a larger particle
size than a more diffuse cloud.
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blogfast25
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Nice treatise on pyrophory, WF...
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DerAlte
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Great posts, all, with special kudos to Annaandherdad for an excellent example of amateur chemistry at its best. If one tries, it is surprising what
can be dome with the simplest of apparatus. The idea of the lighted pinhole to show hydrogen flow is especially innovative! As for mining your own
magnetite from Desert Valley sand, priceless!
Annaandherdad wrote:
| Quote: | | Reduction of CuO by H2 is a demo given in introductory chem classes |
Sadly, I suspect this is not the case. In the USA at least. When I was at school (1950s, UK) learning chemistry every lesson had at least one
demonstration but I have never seen or performed the reductions by hydrogen you describe.
When my eldest son went to his school here in the US, already eager for the subject, he told me that no demonstrations were done. So we did them at
home instead – he already had a Gilbert chemistry set and we amassed some simple apparatus. Soon he did one expt. per day. He aced AP chemistry as a
result. Some sixteen years later #2 son (same school) seemed to manage to escape learning any science that I could perceive. I tried a bit with home
stuff, but all he wanted was bangs and playing computerized games…
Anna is a very lucky girl! If we had a post of the month award, Anna’s Dad would get my vote.
@ Blogfast – I was about to look up the thermochemistry but you saved me the trouble. As usual!
Regards,
Der Alte
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annaandherdad
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blogfast---I recall seeing some posts on this forum with enthaply calculations. Was it you? I taught the UD physics course once on stat mech, but
couldn't follow what you said in that other post. But just for now, do you mind if I ask you some things about the numbers you quote? Like for
starters, is HoF "heat of formation"? And what does it mean, eg in the case of Fe3O4, the heat of formation of Fe3O4 starting from Fe and O2? I'm
just guessing, thanks.
WF, thanks for your suggestions, let me get back to you a little later.
Der Alte, you are much too kind. There are demos of the CuO+H2 on youtube, one from Brazil, but they don't show the "burn" through the pile of CuO
that we saw, at least not the demos I've seen. I took chemistry in high school but never in college, I just assumed they did such demos in college
classes. Your two sons may have differed in temperament as well as date of birth, it could explain their different responses to the science you
offered. But it is a singular pleasure to have a child who is curious, about anything, not just science.
I share with many folks on this forum the sense of loss of a former time when one could just do experiments without all the difficulties thrown in
the way today. I didn't go buy cyanide from the drugstore like Uncle Tungsten, but I did buy acids, mercury and other stuff from chemical mail order
houses, with no trouble. I mowed lawns to make the money for some of my purchases. Now at a more advanced age I am lucky enough to relive some of
the excitement of my teenage years with my daughter. The best part is the thrill of actually doing something and watching it happen, especially when
it's something one was not expecting and has never seen before.
Any other SF Bay chemists?
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blogfast25
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| Quote: Originally posted by annaandherdad | blogfast---I recall seeing some posts on this forum with enthaply calculations. Was it you? I taught the UD physics course once on stat mech, but
couldn't follow what you said in that other post. But just for now, do you mind if I ask you some things about the numbers you quote? Like for
starters, is HoF "heat of formation"? And what does it mean, eg in the case of Fe3O4, the heat of formation of Fe3O4 starting from Fe and O2? I'm
just guessing, thanks.
|
Could have been me but I’m not the only one doing thermocalcs here.
Yes, ‘HoF’ stands for Standard Heat of Formation (or Standard Enthalpy of Formation), the enthalpy that would be released by forming a compound
from its constituent elements, in STP conditions.
Like all reaction heats it is a State Function: e.g. the HoF of say NaCl:
Na(s) +1/2 Cl2(g) === > NaCl(s)
… is independent of how you arrive from State 1 (left hand side) to State 2 (right hand side), it only depends on State 1 and State 2, with ΔH.
= H2 – H1. This statement is also known as Hess Law.
In many cases the HoF of a compound is something that cannot be determined directly from measurement because reactions forming compounds with three or
more elements are rarely practical, e.g.:
K(s) + 1/2 Cl2(g) +3/2 O2(g) === > KClO3(s)
… is highly unlikely to proceed, but with Hess Law and a bit of chemical juggling we can determine its HoF nonetheless. E.g. consider the
decomposition of KClO3 on gentle heating:
KClO3(s) === > KCl(s) + 3/2 O2(g) (R.1)
Assume we measure calorifically the enthalpy of that reaction in STP and call it ΔH.
Now re-write the reaction as the sum of two:
KClO3(s) === > K(s) + 1/2 Cl2(g) +3/2 O2(g) (R. 2)
And:
K(s) +1/2 Cl2(g) === >KCl(s) (R.3)
For R.2 the Heat of Reaction is the opposite of KClO3’s Heat of Formation, call it – ΔH<Sub>F, KClO3 </sub> and for R.3 the Heat
of Reaction is the Heat of Formation of KCl (a known value), call it ΔH<Sub>F, KCl </sub>.
Add R.2 to R.3 and you obtain R.1 and with Hess Law:
– ΔH<Sub>F, KClO3 </sub> + ΔH<Sub>F, KCl </sub> = ΔH
Or:
ΔH<Sub>F, KClO3 </sub> = ΔH<Sub>F, KCl </sub> - ΔH
Et voila, we’ve just indirectly determined the HoF of KClO3.
This method can be applied to countless cases where direct measurement is impossible, by 'creating' a reaction path that is made up of reactions with
known HoRs. Then exploit Hess and 'hey presto!'
Welcome to the world of thermochemistry!
[Edited on 25-4-2012 by blogfast25]
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