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AJKOER
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Neil: My suggestions also included the HCl/H2O2 route and as you are recommending "a beaker full of HCl digesting a bit of iron", I guess you are
endorsing that path.
I agree that the oil drum approach is "off the scale". Note, via dilute HOCl the gas generation is light, and I do my small scale generation in a
closed plastic vessel with the occasional gas release. This also encourages Fe++ and Cl2 interaction without heating. The unreactive gas is possibly
H2. As the reaction proceeds in dilute solutions, I suspect the reaction itself is multi-step and complex (I will spare you all the details which I
have included in another Sciencemadness thread). The Ferric chloride formed is nothing you want to smell (strong Chlorine scent) so the closed vessel
approach is a plus, but still perform outdoors.
NH4OH is safe and effective in neutralizing to obtain Fe(OH)3. When I did this on a very old solution (intense reddish brown) in which I leave Fe
metal, which have become very concentrated as water is consumed in the reaction, most likely FeCl3 and/or Ferric Acetate at that point, the Fe(OH)3
product resembled maple syrup. Normally, pure FeCl3 in water in time will just hydrolyze and leave a deposit of Fe2O3 in a clear aqueous HCl solution.
The off the scale issue is why I strongly suggest a testing and planning approach.
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Neil
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Scale, in large batches the fumes get very bad. I remember someone needing to make a pile of rust and trying bleach + vinegar in a large batch inside
a building. Lots and lots of fumes.
You could drop H2O2 in or dollop of nitric but why, just wait a little longer - the results are the same with less cost.
Iron wants to rust, it really wants to rust. If you are in a massive hurry just set up a charcoal furnace and set some cast iron on fire - grind the
resulting chunks of iron oxide and dry roast.
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weiming1998
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I tried my vinegar approach to make iron rust.
200mls of cleaning vinegar is put in a glass beaker. Iron nails is thrown in. The mix is brought to a steady boil. When it is boiled down to about 50
mls of brown liquid, I stopped heating it, cooled it down, poured 200mls of water in it and transferred it into another container. The remainder of
iron nails was put away. NaHCO3 is poured in. Soon a small layer of brown FeCO3 was precipated out. Excess water is poured off and it is put on a
stove to boil most of the water away. Now I took the FeCO3 out with a spoon and puts it on a metal plate on a stove. This completely dries it and
drives the CO2 away. Now, when the plate is cooled down, a metal knife is used to scrape away the iron oxide. I have got about 20-30 grams of red
Fe2O3. If I used more vinegar, and brought it to a slow boil with aluminum foil covering the top of the beaker, then the yield would have increased.
This seems to be a pretty good method for people with a lack of chemicals, considering that it is much faster than electrolysis with iron electrode.
Also, an aluminum pan can probably handle the vinegar, as I tested the corrosiveness of it with a small strip of aluminum foil. after about 20 minutes
of it in boiling vinegar, it doesn't even look a bit corroded.
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AJKOER
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Per this reference, heating FeCO3 in O2 is more complex in its possible products (Fe3O4 and Fe2O3 as maghemite).
For example, with a trace of O2 and high CO2 pressure at 360 C:
3 FeCO3 + 1/2 O2 --> Fe3O4 + 3 CO2
http://www.scribd.com/fmajdnia/d/74170110/82-Thermal-decompo...
Another approach for those who prefer to work with Iron(II) carbonate, dissolve FeCO3 in carbonated water forming Iron hydrogencarbonate, Fe(HCO3)2.
In air with time, this reputedly decomposes into Fe2O3.xH2O.
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weiming1998
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Quote: Originally posted by AJKOER  | Per this reference, heating FeCO3 in O2 is more complex in its possible products (Fe3O4 and Fe2O3 as maghemite).
For example, with a trace of O2 and high CO2 pressure at 360 C:
3 FeCO3 + 1/2 O2 --> Fe3O4 + 3 CO2
http://www.scribd.com/fmajdnia/d/74170110/82-Thermal-decompo...
Another approach for those who prefer to work with Iron(II) carbonate, dissolve FeCO3 in carbonated water forming Iron hydrogencarbonate, Fe(HCO3)2.
In air with time, this reputedly decomposes into Fe2O3.xH2O.
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Sodium hydroxide (even impure drain cleaner ones) will probably work in creating Fe(OH)3. I was just trying to use as easy to get chemicals as
possible.
By the way, iron can exist as a bicarbonate?
If it can exist as a bicarbonate, then wouldn't the addition of NaHCO3 make iron bicarbonate in the first place? If this is true, then it will be
easier! This would be a very convenient method of people creating iron oxide in countries much more strict on chemicals than Western Australia, using
just baking soda, vinegar and iron nails!
[Edited on 15-2-2012 by weiming1998]
[Edited on 15-2-2012 by weiming1998]
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AJKOER
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Per the attached paper (an interesting fuel cell/ CO2 sequestration discussion):
Fe + 2 CO2 + 2 H2O ---> Fe(HCO3)2 + H2 (g) [1]
Fe(HCO3)2 ---Elevated pH--> FeCO3 (s) + CO2 (g) + H2O [2]
Note, my prior source indicated in the presence of O2 the creation of Fe2O3.xH2O, but reaction [2] is intended to be underground forming instead
FeCO3.
LINK:
http://www.anl.gov/PCS/acsfuel/preprint%20archive/Files/49_1...
Another source on Fe(HCO3)2:
"Iron occurs in many forms in natural water supplies. The most common forms are described below.
1. DISSOLVED IRON: Ferrous bicarbonate [Fe(HCO3)2] is found only in oxygen free water. Dissolved iron is measured in parts per million (ppm). One ppm
is equivalent to approximately 1/4 ounce of iron in 1,900 gallons of water. The recommended limit of iron in drinking water is 0.3 ppm and will begin
staining at 0.5 ppm. The water containing it is clear and colorless when drawn. Upon contact with the air, oxygen is absorbed and reacts with the
dissolved iron to form insoluble ferric hydroxide (commonly known as rust). This clouds the water and colors it in shades of yellow to red-brown.
This reaction produces carbon dioxide as follows
2Fe(HCO3)2 + 1/2O2 + H2O = Fe(OH)3 + 4CO2"
LINK:
http://www.waterwell.cc/IRON.HTM
Now, if one were to react an Iron salt with NaHCO3 (a very basic salt) in excess in an oxygen limited environment, then per reaction [2], one gets a
precipitate of chiefly FeCO3, but if NaHCO3 is not in excess in the presence of CO2 and an acid Fe salt, I would suspect some soluble Fe(HCO3)2 could
be formed. Any O2 in the vessel would also add Fe2O3.xH2O.
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LanthanumK
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I doubt that iron(II) bicarbonate would be anything more than an intermediate in a reaction.
hibernating...
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blogfast25
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Making Fe2O3 from ‘things lying around the house’ is a bit of a rite of passage, one that most grow out of quickly (a bit like MnO2 from battery
electrolyte). Fe2O3 is so cheaply available that it’s not really worth doing at home: even with vinegar and bleach you’re likely to spend more to
make it than to buy it.
| Quote: Originally posted by AJKOER | Per this reference, heating FeCO3 in O2 is more complex in its possible products (Fe3O4 and Fe2O3 as maghemite).
For example, with a trace of O2 and high CO2 pressure at 360 C:
3 FeCO3 + 1/2 O2 --> Fe3O4 + 3 CO2
http://www.scribd.com/fmajdnia/d/74170110/82-Thermal-decompo...
Another approach for those who prefer to work with Iron(II) carbonate, dissolve FeCO3 in carbonated water forming Iron hydrogencarbonate, Fe(HCO3)2.
In air with time, this reputedly decomposes into Fe2O3.xH2O.
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In STP conditions, no matter what insoluble Fe (II) stuff your start from, you end up with Fe2O3. Your reference is describing forcing conditions that
no one here will encounter.
As regards AJoker’s (a bit higher up):
NaOCl + CH3COOH --> HOCl + NaCH3COO
That’s not even a chemical reaction. The hypochlorite solution is completely dissociated into Na+ and ClO-. Acetic acid is a weak acid, in vinegar
dissociated to about 1 % (HAc(aq) + H2O(l) < === > H3O+(aq) + Ac(-))
Hypochloric acid is unstable but a strong acid nonetheless. What happens when you mix the solution of a salt of a strong acid (NaClO) with the
solution of a weak acid? No prizes for guessing: N-O-T-H-I-N-G!
Such a mixture would probably dissolve iron slowly, with the H3O+ oxidising the Fe to Fe2+:
Fe === > Fe2+ + 2 e-
2 H3O+ + 2 e- === > H2 + 2 H2O
And the hypochlorite then oxidising the Fe (II) to Fe (III). Fe(OH)3 would probably precipitate immediately because it’s so insoluble (solubility
product K<sub>s</sub> in the order of 10<sup>-35</sup>, if memory serves me well)
The cheapest/fastest route to home made Fe2O3 may be: dissolve iron scrap in battery acid, warmish to hot for speed. Cool down completely and oxidise
with (slowly added) cold hydrogen peroxide (cheap as chips). Just boiling should precipitate out all the Fe(OH)3, due to hydrolysis of the formed
Fe(III) sulphate. Separate somehow and semi-calcine to drive off water.
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blogfast25
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Yes. Iron rich streams and becks contain Fe(HCO3)2. On oxidation and subsequent hydrolysis that forms Fe(OH)3 which gets deposited on the river bed.
FeCO3 can be precipitated from a neutral Fe(II) solution with sodium carbonate. In air it rapidly oxidises to Fe(III).
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AJKOER
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Per the attached paper (an interesting fuel cell/ CO2 sequestration discussion):
Fe + 2 CO2 + 2 H2O ---> Fe(HCO3)2 + H2 (g) [1]
Fe(HCO3)2 ---Elevated pH--> FeCO3 (s) + CO2 (g) + H2O [2]
Note, my prior source indicated in the presence of O2 the creation of Fe2O3.xH2O, but reaction [2] is intended to be underground forming instead
FeCO3.
LINK:
http://www.anl.gov/PCS/acsfuel/preprint%20archive/Files/49_1...
Another source on Fe(HCO3)2:
"Iron occurs in many forms in natural water supplies. The most common forms are described below.
1. DISSOLVED IRON: Ferrous bicarbonate [Fe(HCO3)2] is found only in oxygen free water. Dissolved iron is measured in parts per million (ppm). One ppm
is equivalent to approximately 1/4 ounce of iron in 1,900 gallons of water. The recommended limit of iron in drinking water is 0.3 ppm and will begin
staining at 0.5 ppm. The water containing it is clear and colorless when drawn. Upon contact with the air, oxygen is absorbed and reacts with the
dissolved iron to form insoluble ferric hydroxide (commonly known as rust). This clouds the water and colors it in shades of yellow to red-brown.
This reaction produces carbon dioxide as follows
2Fe(HCO3)2 + 1/2O2 + H2O = Fe(OH)3 + 4CO2"
LINK:
http://www.waterwell.cc/IRON.HTM
Now, if one were to react an Iron salt with NaHCO3 (a very basic salt) in excess in an oxygen limited environment, then per reaction [2], one gets a
precipitate of chiefly FeCO3, but if NaHCO3 is not in excess in the presence of CO2 and an acid Fe salt, some soluble Fe(HCO3)2 could be formed. Any
O2 in the vessel would also add Fe2O3.xH2O.
These reactions suggest a slow but simple ingredients way to make Fe2O3. Add Fe source to the base of a triangle flash along with pure seltzer water
(H2CO3). Line the base of this large flash with tubing with holes connected to an air pump (common supplies for those with a fish tank). Periodically
refresh the H2CO3 as needed. The logic is to create Fe(HCO3)2 per reaction [1] and add O2 to form Fe2O3.xH2O and CO2. Now, per the vessel's design,
some of the newly created heavy CO2 gas may, hopefully, be captured and recycled per [1]. Caution, run this experiment in a ventilated area as
Hydrogen has a lower explosion limit (or LEL) of 4% (measured as % of volume in air, see http://www.delphian.com/chc.htm ). Of course, if you live in a heavily polluted area now, just leave your iron scraps outside. Those still
enjoying fresh air and near trees may wish to run this experiment.
[Edited on 15-2-2012 by AJKOER]
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AJKOER
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| Quote: Originally posted by blogfast25 |
As regards AJoker’s (a bit higher up):
NaOCl + CH3COOH --> HOCl + NaCH3COO
That’s not even a chemical reaction. The hypochlorite solution is completely dissociated into Na+ and ClO-. Acetic acid is a weak acid, in vinegar
dissociated to about 1 % (HAc(aq) + H2O(l) < === > H3O+(aq) + Ac(-))
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Hi Blogface:
Forming HOCl from NaOCl and almost any weak acid (Boric, Carbonic, Ascorbic, Acetic, or any very dilute mineral acid) is best represented, in my
opinion, by the net ionic equation:
OCl(-) + H(+) ---> HOCl
as Hypochlorous acid is such a weak acid (lack of ionization) that one author suggests just write it as HOCl. EDIT: Actually, this is no longer my
opinion alone as I have found an examine question and the above net ionic equation is cited precisely as the correct answer; see question 11 at: http://aasiri2.kau.edu.sa/Files/0002617/Files/28029_Chapter1...
This fact is taking to advantage in the preparation of Hypochlorous acid by the action of a metal oxide (like HgO or ZnO) or select carbonates on
Chlorine water:
Cl2 + H2O <----> HOCl + H(+) + Cl(-)
as the metal oxides much more rapidly attack and remove the HCl. This leaves a weak solution of HOCl and a metal salt which can be distilled to
produce HOCl. Alternately, one could employ a solvent per Patent 3718598 (solvents include acetone, methyl ethyl ketone, methyl isobutyl ketone,
diethyl ketone, di-n-propyl ketone, methyl cyanide, ethyl cyanide, methyl acetate, ethyl acetate and methyl propionate) to extract a nearly chloride
free HOCl.
[Edited on 16-2-2012 by AJKOER]
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Neil
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The fastest and the only way to make Fe2O3 at home for less then you can buy it, is to set the Iron on fire.
Iron metal costs less gram per gram then iron oxides unless you get it in bulk (train car sized loads).
But: as soon as you start trying to use chemicals you are going to be paying more for it then you would from a pottery supply place or by just going
out and looking or a bog and then picking up big chunks of iron/oxide/carbonate for free that stuff only needs roasting and grinding.
Tis not complicated, the efforts forthwith notwithstanding.
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AJKOER
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| Quote: Originally posted by blogfast25 |
The cheapest/fastest route to home made Fe2O3 may be: dissolve iron scrap in battery acid, warmish to hot for speed. Cool down completely and oxidise
with (slowly added) cold hydrogen peroxide (cheap as chips). Just boiling should precipitate out all the Fe(OH)3, due to hydrolysis of the formed
Fe(III) sulphate. Separate somehow and semi-calcine to drive off water.
|
Blogfast25:
You are definitely partially right. I agree that:
Fe + H2SO4 --> FeSO4 + H2 (g)
and
4 FeSO4 + O2 + 2 H2SO4 --> 2 Fe2(SO4)3 + 2 H2O
where you are using H2O2 in place of O2. However,
Fe2(SO4)3 + H2O --> Fe2(SO4)O + H2SO4
Reference: Patent 3078180
which implies to me that the hydrolysis of Ferric Sulphate may produce a basic Ferric Sulphate, and not Fe2O3.xH2O immediately as you claimed.
So don't rush to perform this synthesis, it may not work!
[Edited on 16-2-2012 by AJKOER]
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weiming1998
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| Quote: Originally posted by Neil | The fastest and the only way to make Fe2O3 at home for less then you can buy it, is to set the Iron on fire.
Iron metal costs less gram per gram then iron oxides unless you get it in bulk (train car sized loads).
But: as soon as you start trying to use chemicals you are going to be paying more for it then you would from a pottery supply place or by just going
out and looking or a bog and then picking up big chunks of iron/oxide/carbonate for free that stuff only needs roasting and grinding.
Tis not complicated, the efforts forthwith notwithstanding.
|
How will you set iron on fire? Using just a crucible, iron and furnace? That method is cheap, but takes a very long time with full heat in a furnace.
If you have an electric furnace that can melt iron, then it will be as fast as oxidizing carbon, but very few people have electric furnaces. Also,
that creates Fe3O4, not the Fe2O3 that we are looking for.
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blogfast25
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| Quote: Originally posted by AJKOER |
Hi Blogface:
Forming HOCl from NaOCl and almost any weak acid (Boric, Carbonic, Ascorbic, Acetic, or any very dilute mineral acid) is best represented, in my
opinion, by the net ionic equation:
OCl(-) + H(+) ---> HOCl
as Hypochlorous acid is such a weak acid (lack of ionization) that one author suggests just write it as HOCl. |
You’re right that HClO is a weak acid. It changes very little though: the metal is oxidised by H3O+ to Fe (II), then further by the hypochlorite to
Fe (III).
You also need to take into account the relative proportions of [ClO-] and [HAc] in the initial solutions. Commercial vinegar is typically about 0.8 M,
thin commercial bleach about 4 - 5 % NaClO. If there’s excess HAc all the above is quite academic, as the H3O+ will still be mainly supplied by the
vinegar.
| Quote: Originally posted by AJKOER | | Quote: Originally posted by blogfast25 |
The cheapest/fastest route to home made Fe2O3 may be: dissolve iron scrap in battery acid, warmish to hot for speed. Cool down completely and oxidise
with (slowly added) cold hydrogen peroxide (cheap as chips). Just boiling should precipitate out all the Fe(OH)3, due to hydrolysis of the formed
Fe(III) sulphate. Separate somehow and semi-calcine to drive off water.
|
Blogfast25:
However,
Fe2(SO4)3 + H2O --> Fe2(SO4)O + H2SO4
Reference: Patent 3078180
which implies to me that the hydrolysis of Ferric Sulphate may produce a basic Ferric Sulphate, and not Fe2O3.xH2O immediately as you claimed.
So don't rush to perform this synthesis, it may not work!
[Edited on 16-2-2012 by AJKOER] |
The last sentence shows as ever how much your ‘knowledge’ comes from obsessive Interwebs scanning and how little from practical experience.
It is theoretically true that hydrolysis can form a basic ferric sulphate, more likely Fe(OH)SO4 hydrate, which I believe I’ve seen but that Fe(OH)3
can also form and I’ve also seen that form. Which of the two actually forms depends on iron concentration, pH and temperature.
Suppose though that the basic ferric sulphate forms (in practical terms this means that the hydrolysis didn’t go ‘all the way’), all that is
needed is to increase pH, and not by much either:
For Fe3+(aq) + 3 OH-(aq) === > Fe(OH)3(s) the equilibrium constant (solubility product) is in the order of 10<sup>-35</sup>. Depending
on concentration, Fe3+ starts dropping out of solution from about pH 4, not much alkali is needed to achieve that.
But if you’re working with quite concentrated solution ([Fe] > 1 M, as is practical) then chances are high that the hydroxide will drop out upon
heating, w/o any alkali addition, unless of course your solution is mega acidic, in which case one was wasting acid.
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AJKOER
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| Quote: Originally posted by blogfast25 |
| Quote: Originally posted by AJKOER | Per this reference, heating FeCO3 in O2 is more complex in its possible products (Fe3O4 and Fe2O3 as maghemite).
For example, with a trace of O2 and high CO2 pressure at 360 C:
3 FeCO3 + 1/2 O2 --> Fe3O4 + 3 CO2
http://www.scribd.com/fmajdnia/d/74170110/82-Thermal-decompo...
Another approach for those who prefer to work with Iron(II) carbonate, dissolve FeCO3 in carbonated water forming Iron hydrogencarbonate, Fe(HCO3)2.
In air with time, this reputedly decomposes into Fe2O3.xH2O.
|
In STP conditions, no matter what insoluble Fe (II) stuff your start from, you end up with Fe2O3. Your reference is describing forcing conditions that
no one here will encounter.
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Blogfast25:
My statement appears to be factually more correct. Per the work cited below, the product on thermal decomposition of FeCO3 is Fe3O4 and Fe2O3. To
remove the Fe3O4, one must further anneal in pure oxygen for 2 hours at 500 C.
http://gong.ustc.edu.cn/Article/2008A01.pdf
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AJKOER
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| Quote: Originally posted by blogfast25 |
"The cheapest/fastest route to home made Fe2O3 may be: dissolve iron scrap in battery acid, warmish to hot for speed. Cool down completely and oxidise
with (slowly added) cold hydrogen peroxide (cheap as chips). Just boiling should precipitate out all the Fe(OH)3, due to hydrolysis of the formed
Fe(III) sulphate. Separate somehow and semi-calcine to drive off water."
and now:
"It is theoretically true that hydrolysis can form a basic ferric sulphate, more likely Fe(OH)SO4 hydrate, which I believe I’ve seen but that
Fe(OH)3 can also form and I’ve also seen that form. Which of the two actually forms depends on iron concentration, pH and temperature.
Suppose though that the basic ferric sulphate forms (in practical terms this means that the hydrolysis didn’t go ‘all the way’), all that is
needed is to increase pH, and not by much either:"
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So, if I understand your "cheapest/fastest" synthesis correctly, Iron plus battery acid plus H2O2 plus ice to cool down, plus, if all else fails NaOH
(depending on temperature, concentration and pH) and, of course, lots of boiling....followed by "Separate somehow and semi-calcine to drive off
water." Really?
Would anyone else like to comment?
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White Yeti
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Don't jump through hoops, simply let some iron rust away in salt water.
Materials:
-some jars
-some iron/steel
-some water
-some salt
-a couple a months
Procedure:
-add iron to jar
-add water to iron
-add salt to water
-wait a month or so
-collect product, filter and dry
-repeat
Good things come to those who wait. I got ~200g of oxide by this method. No need for that sulfuric acid/peroxide nonsense. I already posted a picture
of my stash here on this thread. I highly recommend this method, and believe me, I tried almost every method there is to make iron oxide, this is the best one so far.
"Ja, Kalzium, das ist alles!" -Otto Loewi
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weiming1998
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| Quote: Originally posted by White Yeti | Don't jump through hoops, simply let some iron rust away in salt water.
Materials:
-some jars
-some iron/steel
-some water
-some salt
-a couple a months
Procedure:
-add iron to jar
-add water to iron
-add salt to water
-wait a month or so
-collect product, filter and dry
-repeat
Good things come to those who wait. I got ~200g of oxide by this method. No need for that sulfuric acid/peroxide nonsense. I already posted a picture
of my stash here on this thread. I highly recommend this method, and believe me, I tried almost every method there is to make iron oxide, this is the best one so far.
|
Takes a very, very long time though. Most people won't have the patience.
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Neil
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| Quote: Originally posted by weiming1998 |
How will you set iron on fire? Using just a crucible, iron and furnace? That method is cheap, but takes a very long time with full heat in a furnace.
If you have an electric furnace that can melt iron, then it will be as fast as oxidizing carbon, but very few people have electric furnaces. Also,
that creates Fe3O4, not the Fe2O3 that we are looking for. |
Dig a hole in the ground, line it with clay or just pack the soil well. lay down a steel pipe and attach a air bower to the pipe. fill hole with wood
or charcoal and set it on fire, tun on air. when inside of heap is white hot start feeding steel in, sizzling heat
is easily reached with charcoal and air, that is the white hot temperature where pieces of steel make the sound of bacon on a hot grill and spit
sparks like a toy sparkler aka they are on fire.
Keep adding fuel and steel making sure that you are destroying all the metal you add and that it is not just melting and pooling at the end take the
slag grind it up boil it dry it and separate everything with a magnet.
Fe3O4 is rather soluble in silica flux while Fe2O3 precipitates out under oxidizing conditions, keep lots of air moving into the burning pile but not
so much that you are blowing the heat out of the fire.
Cast Iron burns the best, a hot air blast on it makes it swell up like a sponge absorbing water until it is a porous briquette of iron oxide. Too
little oxygen in the blast and you just end up with molten metal.
It creates FeO, Fe2O3 and Fe3O4. Grinding and Roasting will convert a lot of the matter to Fe2O3. The black powder that results grinds reddish and is
not very magnetic.
If the ground wasn't frozen and under a lay of ice I'd take a picture for ya.
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White Yeti
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You're right, but this "synthesis" is care free, you can get it to work without supervision or any kind of intervention. If you get several units to
work in parallel, you can get quite a large amount. This method is working very well for me, I'm getting about 10 grams of oxide every month.
I used to electrolyse water with iron nails and I got ~5g at best after HOURS of supervision. The choice is yours.
Besides, it's not like you need to make thermite very often. If you want large amounts of iron oxide, just buy it.
"Ja, Kalzium, das ist alles!" -Otto Loewi
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weiming1998
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| Quote: Originally posted by White Yeti |
You're right, but this "synthesis" is care free, you can get it to work without supervision or any kind of intervention. If you get several units to
work in parallel, you can get quite a large amount. This method is working very well for me, I'm getting about 10 grams of oxide every month.
I used to electrolyse water with iron nails and I got ~5g at best after HOURS of supervision. The choice is yours.
Besides, it's not like you need to make thermite very often. If you want large amounts of iron oxide, just buy it. |
Ok then, I guess that's the cheapest route, requiring only salt, iron and air.
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AJKOER
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OK, salt, water and air. But, to add some chemistry, for example, NaCl allows the pitting of Aluminum as it dissolves the protective Al2O3 by raising
the pH followed by the hydrolysis of Al in H2O:
2 Al + 6 H2O ---> 2 Al(OH)3 + 3 H2 (g)
Now, in the case of Iron, the elevated pH help dissolves the protective FeO/F2O3 coating and CO2 (from the air or dissolved in the tap water) attacks
the Iron:
Fe + 2 CO2 + 2 H2O ---> Fe(HCO3)2 + H2 (g)
and with air:
2 Fe(HCO3)2 + 1/2 O2 + H2O ---> Fe(OH)3 + 4 CO2
the rust forms. However, be careful to not raise the pH via NaCl with limited air flow as you will also form Iron carbonate:
Fe(HCO3)2 ---Elevated pH--> FeCO3 (s) + CO2 (g) + H2O
So I would recommend, add some salt to carbonated water in an air flow (like fish tank air tubing) and the process should be more efficient.
[Edited on 20-2-2012 by AJKOER]
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Neil
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| Quote: Originally posted by AJKOER | OK, salt, water and air. But, to add some chemistry, for example, NaCl allows the pitting of Aluminum as it dissolves the protective Al2O3 by raising
the pH followed by the hydrolysis of Al in H2O:
2 Al + 6 H2O ---> 2 Al(OH)3 + 3 H2 (g)
Now, in the case of Iron, it is CO2 dissolving the Iron:
Fe + 2 CO2 + 2 H2O ---> Fe(HCO3)2 + H2 (g)
and with air:
2 Fe(HCO3)2 + 1/2 O2 + H2O ---> Fe(OH)3 + 4 CO2
the rust forms. However, be careful to not raise the pH via NaCl with limited air flow as you will also form Iron carbonate:
Fe(HCO3)2 ---Elevated pH--> FeCO3 (s) + CO2 (g) + H2O [2]
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?
NaCl solutioms do not dissolve Al2O3. The Cl is able to penetrate the Oxide layer to allow the Al underneath to react via electrochemical oxidistion
just like the pitting of Iron by chloride ions. Molten NaCl will flux Al2O3, is that what you are thinking of?
Attachment: Pitting of alumina by Chloride ions.pdf (216kB)
This file has been downloaded 5143 times
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White Yeti
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| Quote: Originally posted by AJKOER | OK, salt, water and air. But, to add some chemistry, for example, NaCl allows the pitting of Aluminum as it dissolves the protective Al2O3 by raising
the pH followed by the hydrolysis of Al in H2O:
2 Al + 6 H2O ---> 2 Al(OH)3 + 3 H2 (g)
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Double facepalm:
Since when is aluminium metal corroded by salt water?! [sarcasm] I'm sure people who use anodised aluminium in corrosive environments are just a bunch
of pansies. After all aluminium exhibits pitting corrosion in salt water. As we all know, NaCl is a notorious base, 'tis basic enough to eat through
the aluminium oxide passivation layer.[/sarcasm off]
"Ja, Kalzium, das ist alles!" -Otto Loewi
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