Sciencemadness Discussion Board - Iron Oxide for use in Thermite

An inexpensive route but you will have to experiment on small scale to test whether the process is capable of producing sufficient quantity for your needs. Procedure is:

1. Mix vinegar and Bleach to make Hypochlorous acid in a solution with Sodium acetate. Use excess of vinegar (Acetic acid) as the NaOCl can vary from 3% to 6% in strength. Reaction:

NaOCl + CH3COOH --> HOCl + NaCH3COO

2. Add Iron (solid sheet, scraps or filings). There will be some bubbling as dilute HOCl is capable of attacking iron. With time, the solution becomes reddish brown with a ferric salt (FeCl3 and some Ferric acetate).

3. Add a convenient base (NaOH or NH4OH) to precipitate Fe(OH)3. Filter, wash and dry to obtain Fe2O3.

To speed up the process (careful, the reaction can get out of control bubbling Chlorine), use HCl + H2O2 (creates concentrated Chlorine water) in place of the HOCl preparation as:

2 HCl + O2 ---Heat--> 2 HOCl

is one of HCl + H2O2 reaction products although with the decomposition of H2O2 releasing oxygen. Continue with Steps 2 and 3.

There are YouTube videos on the reaction of Iron and HCl/H2O2. The third metal shown here, for example, is Fe.

http://www.youtube.com/watch?feature=endscreen&NR=1&...

Good luck.

Why bleach and vinegar, then NaOH to precipate? Why not just buy FeSO4 and add NaOH? If you seriously lack chemicals, why not just heat up some vinegar/HCl with chunks of iron in it, then precipate with NaOH?

No offense, but all your chemistry reactions seems to involve either chlorine or hypochlorites. What's the fascination with hypochlorites?

[Edited on 28-1-2012 by weiming1998]

AJKOER - 28-1-2012 at 13:11

If this person more likely has large amount of FeSO4 available, I agree.

As Bleach is available in food stores, it is cheap and strong oxidizer, my 1st choice. Note, Bleach is not necessarily the oxidizer employed, it could be a derivative (like Chlorine water, HCl, Cl2, HOCl, Cl2O or even Ferrates). The Bleach/Acetic acid combination to HOCl is quite interesting as the Sodium acetate is an apparent catalyst in some instances. Also, storing HOCl does not work well decomposing into weak HCl and producing O2 gas, or worst some HClO3 with diffused light to add chlorates.

I also like H2O2/ferrous salt and percarbonates (H2O2 with the catalyst Na2CO3 plus heat).

I try to avoid storing strong acids around as I have kids into everything, and throwing things that can't open.

Also, it has gotten to the point in America where anybody having chemicals or glassware is presumed to be......and not a garage chemist. Having served on a jury many times, I can attest to the fact that jurists are complete bias idiots no matter what the facts presented, so that is yet another reason not to have stuff around.

Neil - 28-1-2012 at 14:24

DO SOME EXPERIMENTING.

This is 100% useless to produce large amounts of Iron oxide, try it you will see.

"An inexpensive route but you will have to experiment on small scale to test whether the process is capable of producing sufficient quantity for your needs. Procedure is:"

An inexpensive route that you will have to test to see if it is truly inexpensive or even useful?

There is no way that dissolving Iron with bleach will come close to being as cheap as buying it from a pottery supply store or Ebay.

Bleach is not cheap, the water in bleach is cheap but saying bleach is cheap is like saying acetic acid from vinegar is cheap. I have tried making large amounts of Fe2O3 this way, it is just a waste of resources.

@weiming: FeO is unstable and will happily convert to Fe2O3 given the chance. Pure powdered FeO is rumoured to be pyrophoric.

@ManBearSwine: Yes, Your right I'm sorry I did not realise that you were already heating to the needed degree.

AJKOER - 28-1-2012 at 16:53

Neil: Yes, I agree that the dilute HOCl may not be a fast or cheap path to make Fe2O3. With respect to yield, per my suggested testing, you may have come to realize you need a oil drum of bleach/vinegar (cleaning supply store may provide higher strength and larger containers). If this is possible for you, then scale is not an issue (cost perhaps).

To be honest, we have not really factored the cost of labor/time in making fine Fe powder (or not), and of converting FeO to Fe2O3 (if necessary). In fact, if we have fine Iron powder, just adding Bleach may work to form some FeO quickly, but it still has to be further air oxidized. The advantage of my method is that it does dehydrate to Fe2O3 directly, while even strong HCl forms FeCl2, which upon adding a base, leads to Fe(OH)2 (or FeO.H2O).

Now, HCl/H2O2 is much more costly but (see the video) a vigorous way to dissolve solid Iron. Still, I would conduct a test. Weigh my Fe and in 10 minutes, remove and re-weigh the Fe to estimate efficiency/feasibility. Note, the finer the Iron powder (hard work, I have tried filing and surrendered) the better this or any method.

As a matter of record, I have prepared Fe(OH)3 (not Fe2O3) per the dilute HOCl route. It formed an incredible voluminous precipitate and I can see why it is used in water purification. However, I would suspect, if I have dried it, a low yield of Fe2O3, hence my suggested testing.

[Edited on 29-1-2012 by AJKOER]

weiming1998 - 28-1-2012 at 17:07

Actually, dilute acetic acid (vinegar) can actually attack iron pretty vigorously when heated! The iron acetate can then be precipated with NaOH or KOH. Or even precipation with Na2CO3/NaHCO3 (available as washing soda/baking soda), then heat. If FeO is unstable, then it probably will convert to Fe2O3 when cooled (some information suggests that FeO is stable at high temperatures.)

A low temperature simmer of some vinegar in an aluminum pot (I heard that vinegar doesn't attack aluminum because of passivation) with a few chunks of scrap iron, then precipation by OH-/CO3-2/HCO3- and heat should work cheaply and efficiently.

Heated HCl/NaHSO4 does the job better but it can't be contained in anything but a glass container/ The HCl fumes and droplets of NaHSO4 solution goes in the air and irritates people's noses and lungs.

AJKOER - 28-1-2012 at 18:14

Weiming1998: My source indicates that Fe(OH)2 rapidly converts in the presence of air via dark gray-green and black intermediates to Fe2O3.xH2O.

Note, Fe(OH)2 is soluble in acids, and also in hot concentrated NaOH to a slight extent forming hydroxoferrates(II), Na4[Fe(OH)6].

I would not use a carbonate to neutralize as insoluble FeCO3 could be formed. You want Fe(HCO3)2 which requires free CO2 in the water. Iron hydrogencarbonate in air forms Fe2O3.xH2O per my source.

Your info on Aluminum and vinegar is incorrect. Vinegar is reported to weaken the passivation of Al. This is a major concern to the Aluminum foil industry as having your food + vinegar in prolonged contact could result in Al leaching, a major public relations nightmare in light of increasing health concerns associated with Aluminum. I, myself, have pre-soaked Al in vinegar before dissolving in NH4OH (a little quicker).

Weiming1998 your method can similarly be attacked on scale/yield concerns. How big is your pot? (just joking)

[Edited on 29-1-2012 by AJKOER]

weiming1998 - 28-1-2012 at 18:30

I haven't tried a pot yet, I use a glass beaker to contain the vinegar. So aluminum does get attacked by vinegar. Maybe cold HCl/NaHSO4 with Fe might work.. The reaction would be slow. The bleach/ vinegar mixture did yield rust (I tried that before), but I have to wait overnight (which an acid would have corroded the metal away) and the mixture smelt strongly of chlorine, even from a distance. But maybe that is a viable route to people that has a serious lack of chemicals, from regulations, or simply from having no access to various shops. But electrolysis of water by an iron nail is a very inefficient, slow, process. This process certainly doesn't require much chemicals, but costs a lot of electricity.

Neil - 28-1-2012 at 18:44

@AJKOER Just because something can be done - does not make it a good idea, a reasonable idea or a plausible synthesis.

Can you imagine how much fume you would be dealing with if you had a barrel full of bleach and vinegar fizzing away?

Can you imagine how much it would cost to buy that much bleach and vinegar?

Or how long it would take to filter the iron hydroxide sludge out of that much fluid?

Can you honestly suggest that a garbage can full of fizzing acid is a better idea then a beaker full of HCl digesting a bit of iron?

Really, Try things out.

AJKOER - 28-1-2012 at 18:55

OK, to be honest, the reaction of Al and vinegar is very slow (invisible). There is an inception period (meaning that nothing happens and then the reaction slowly commences). So don't change your process if the contact with Al is short-lived.

Using NaCl also is reported to cause pitting on Al metal and has been suggested as a primer in a Sciencemadness thread.

AJKOER - 28-1-2012 at 19:46

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.

Neil - 28-1-2012 at 20:18

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.

weiming1998 - 30-1-2012 at 02:24

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.

AJKOER - 14-2-2012 at 17:55

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.

weiming1998 - 15-2-2012 at 00:19

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.

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]

AJKOER - 15-2-2012 at 06:29

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.

LanthanumK - 15-2-2012 at 06:48

I doubt that iron(II) bicarbonate would be anything more than an intermediate in a reaction.

blogfast25 - 15-2-2012 at 07:00

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.

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 Ks in the order of 10-35, 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.

blogfast25 - 15-2-2012 at 07:04

By the way, iron can exist as a bicarbonate?
[Edited on 15-2-2012 by weiming1998]

[Edited on 15-2-2012 by weiming1998]

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).

AJKOER - 15-2-2012 at 07:31

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]

AJKOER - 15-2-2012 at 11:23

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(-))

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]

Neil - 15-2-2012 at 18:04

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.

AJKOER - 15-2-2012 at 20:19

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]

weiming1998 - 16-2-2012 at 00:50

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.

blogfast25 - 16-2-2012 at 07:05

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.

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 10-35. 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.

AJKOER - 16-2-2012 at 07:27

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.

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

AJKOER - 16-2-2012 at 12:44

"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:"

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?

White Yeti - 16-2-2012 at 17:53

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.

weiming1998 - 17-2-2012 at 01:28

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.

Neil - 17-2-2012 at 06:54

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.

White Yeti - 17-2-2012 at 09:35

Takes a very, very long time though. Most people won't have the patience.

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.

weiming1998 - 17-2-2012 at 16:01

Quote: Originally posted by White Yeti

Takes a very, very long time though. Most people won't have the patience.

Ok then, I guess that's the cheapest route, requiring only salt, iron and air.

AJKOER - 20-2-2012 at 07:54

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]

Neil - 20-2-2012 at 08:18

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]

?

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 4938 times

White Yeti - 20-2-2012 at 10:14

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]

AJKOER - 20-2-2012 at 12:45

Neil:

Here is a paper you once referenced me too. My favorite quote from the first page is "Aluminum and Al alloys are passivated in neutral solutions, so they have better corrosion resistance. However, the passivation film over Aluminum and Al alloys is easily destroyed in solutions containing active anions, such as Cl"

LINK:
http://proj3.sinica.edu.tw/~chem/servxx6/files/paper_7425_12...

So I am guilty of replacing the word "destroyed" with "dissolved", and technically speaking, you are most likely correct. Good point, as it helps explain the actual process.

I say most likely because in the case of a very acidic (or basic) salt, the pH may be sufficient to actually dissolve any dislodged Al2O3, as Alumina, per my source, does dissolve in a strong acid or base. As an example of the latter, a known serious incompatibility is reported to exist between Aluminum metal and Ferric (or Ammonium) Sulphate in the presence of moisture.

[Edited on 20-2-2012 by AJKOER]

White Yeti - 20-2-2012 at 13:24

[...]it dissolves the protective Al2O3 by raising the pH[...]

Nowhere in the paper did the author mention sodium chloride as a strong acid or base.

AJKOER - 20-2-2012 at 14:47

White Yeti:

I see your point that someone may incorrectly infer that I meant that NaCl is such a salt. As to how to best describe a NaCl solution, please see the discussion at this link:

http://www.chemicalforums.com/index.php?topic=19886.0

I have since edited my comment with examples of highly acidic salts.

Neil - 20-2-2012 at 17:36

http://en.wikipedia.org/wiki/Credential

For more on the pH of NaCl solutions

Attachment: Effect of NaCl and HCl on pH.pdf (66kB)
This file has been downloaded 724 times

More on the pH of NaCl under CO2
Attachment: measuring NaCl pH.pdf (34kB)
This file has been downloaded 801 times

It is indeed funny how rare it is for people on this forum to admit to errors.

weiming1998 - 21-2-2012 at 01:31

Quote: Originally posted by Neil

Just what I think about how salt water attacks aluminum slowly over time:

1 mole of salt solution disassociates into 10^-7 moles of both
HCl and NaOH. They react, but more soon disassociates. When you put aluminum in, they both attack the Al2O3, albeit very slowly, forming AlCl3(aq) and NaAlO2/some Al(OH)4 ions. As the oxide layer is very thin, it is soon destructed by the trace HCl and NaOH, and reacts with the water+the trace HCl/NaOH. The reason why NaAlO2 isn't formed is because the HCl reacts with it in water, forming more NaCl and Al(OH)3. The AlCl3 is reacted with the NaOH and similarly, forms the same products as above. At the end, what you are left with are flakes of Al(OH)3.

Neil - 21-2-2012 at 07:09

No, they do not touch the Al2O3. It remains and the corrosion form blisters underneath it. The chloride ions set up a cycle of moving through the Al2O3 barrier and attacking the Al underneath. The swelling of the material is able to break and blister the outer skin of Al2O3 but it sure as heck 100% is not dissolving it to form aluminates and the like. Rubies and Saphire do not dissolve in strong or weak NaCl solutions.

Chloride penetrates the barrier it does not dissolve, or destroy the barrier.

elementcollector1 - 17-6-2013 at 18:11

I've had an odd experience with attempting to make red iron oxide.
(Unbalanced) reactions: Fe + H2SO4 -> Fe2(SO4)3 + H2
Fe2(SO4)3 + NaOH -> Na2SO4 + Fe(OH)3
Fe(OH)3 -> Fe2O3 + H2O

A few problems occured after step 2:
1) Filtering. This is a common problem for ferric hydroxide, as I've heard.
2) As the filter cake was drying, needle-thin crystals appeared, quickly drying the rest out. I assumed these were sodium sulfate and washed and refiltered the precipitate (which was by now easy to filter because it was so dry and clumpy)
3) The filter cake is now a reddish-black powder, and upon heating with a (bad) hotplate, there is no color change or loss of mass. I can't use the blowtorch because it's broken for some reason, and given that I am still living with my family, the oven is obviously not an option. Ideas?
The powder does look like it might be ferric hydroxide, and despite the dark color is not magnetite (no attraction whatsoever to a stirbar magnet).

blogfast25 - 18-6-2013 at 03:59

A couple of points.

Fe is NOT oxidised to ferric sulphate by H2SO4, only to ferrous sulphate (FeSO4). H₃O⁺ === > H₂ just doesn't have that kind of oxidising power.

But by now all Fe will have air oxidised to Fe2O3, so no real worries there.

Fe(OH)3 can be a barstool to filter because of possible peptisation. It occurs when the ionic strength of the wash water becomes low. The whole thing can then (sometimes) simply run through the filter in one fell swoop. Very annoying.

Ferric oxide's colour ranges from ocre to almost black, depending on mode of preparation. Colour is a poor marker.

Dry, even semi-calcinate, the product by heating it on top of your kitchen cooker on high heat in a disused steel pan (NOT aluminium, OBVIOUSLY), using an old steel spoon as stirring rod. I've done this and it works well...

[Edited on 18-6-2013 by blogfast25]

Fantasma4500 - 19-6-2013 at 06:37

i have a question...

now we can choose to do electrolysis.. OR using HOCl formed in several different ways or even creating FeSO4 by reacting it with H2SO4

now what im wondering is... why dont you guys just combine it??
when i was obtaining Cr(OH)2 by HCl + steel i usually kickstarted it with electrolysis, its a different thing but still metal etching with corrosive liquids

steelwool in HOCl solution could possible act as a neutral meaning it would take part in the reaction, only thing that must not be allowed to happen is contact with cathode or anode
i obtained 20g Fe2O3 through electrolysis in just water as i remember it, long procedure and more than just messy

ElizabethGreene - 19-6-2013 at 12:41

I have manufactured, found, and bought Iron Oxide.

I've personally manufactured rust two ways.

First, and most satisfying, is to light fine steel wool on fire. Steel wool is painfully expensive for this though. I paid $3.50 for 12 pads of it yesterday, and they couldn't weigh more than an ounce or two. On the other hand they are a convenient source if you can nick your mom's used brillo pads from under the sink and scrub the soap and grease out of them .

Second, I've manufactured it by electrolysis. Two iron electrodes, placed in a mason jar, with an unspecified quantity of salt added, and connected to a 12 volt automotive battery charger. This ran for a week. The first couple of days seemingly nothing happened, but an oily sheen developed on top of the water. On days 3-6, the water turned light green. This was filtered and allowed to dry whereupon the product turned red. The quantity produced was infinitesimal.

I apologize for my lack of detail on this, it was before I started taking proper notes. One lesson learned was to keep your power supply wires OUT of the water or they will disappear a lot faster than your target iron.

I "found" Iron Oxide in my shed where a plastic bin of masonry nails had filled up with water from a leak. Filtering the water gave a bit of nasty impure rust. Drying the nails and then "tumbling" them in a 2 liter plastic bottle produced a fair amount of rust powder.

Finally, I bought Iron Oxide from Mid-South Ceramic Supply for $4.00 per pound for 5 pounds. This company was on my drive home from work and this was by far the easiest way I've ever found to get rust. I found them by searching the internet for "Ceramic Supply near Nashville Tennessee." They also have Tin/Titanium/Aluminum/and Nickel oxides and Barium/Copper/Lithium/Magnesium/Strontium Carbonates. The most unfortunate bit about them is that it appears I'm going to be spending a fair amount of money there.

Back to Aluminum powder for a moment. I inquired about this at a Sherwin-Williams paint store and they were unfamiliar with the substance. Could someone be slightly more specific as to what kind of paint store to find this in?

Thanks,
Elizabeth

blogfast25 - 20-6-2013 at 03:51

Quote: Originally posted by ElizabethGreene

Here in the UK some model [RC planes, RC cars, dolls' houses etc] shops stock it as a filler for metallic paint or die cast resins. You might find it also in some paint-for-crafts shops.

Dr.Bob - 21-6-2013 at 13:57

Back to Aluminum powder for a moment. I inquired about this at a Sherwin-Williams paint store and they were unfamiliar with the substance. Could someone be slightly more specific as to what kind of paint store to find this in?

Years ago, paint stores would mix up, make, or modify paint. Now they get cans of white paint and tint it with a machine that likely is likely smarter than the people operating it. But few places actually mix things together to make paint unless the machine does it for them. So it is much harder to find raw pigments and paint chemicals now in stores. If you can find an OLD paint store, they might have some left, likely hidden on a shelve in back. Ask the oldest person there for help...

And clearly, buying iron oxide is a simple solution or just finding old rusty things and shaking them well. Using nails or other small pieces of iron will rust plenty fast with some water, salt, and maybe bubble some air through. If you scaled that up a bit, you could generate lots of rust for nearly nothing. You might be able to find a hardware store with some old, rusty nails that got wet that are discounted because of that. Letting nature do its thing is a great way to do chemistry.

blogfast25 - 22-6-2013 at 06:28

Quote: Originally posted by Dr.Bob

Almost none, I should think.

Al powder, probably depending on where you are, isn't hard to get though. Amazon and eBay UK both advertise it. Pyro webstores offer more grades than I care for. It's not OTC but it's not problematic either.

AJKOER - 25-10-2013 at 15:15

OK, upon further studies in the electrochemistry of batteries, there is, in effect, an iron air battery (see visuals at https://www.google.com/search?q=iron+air+battery&tbm=isc... ).

As such, in my opinion, a better description of the reaction of Iron and an oxygen source (like, for example, H2O2 in place of O2) is better described as electrochemical process. Assuming an electrochemical reaction similar to the that of Zinc and Oxygen in aqueous NaCl (see http://sci-toys.com/scitoys/scitoys/echem/batteries/batterie... ), the reaction, in part, could proceed as follows with Iron, Aluminum, H2O2 and NaCl :

At the Iron anode:

Fe + 2 OH- ⇒ Fe(OH)2 + 2e- (see Eq [1] at http://jes.ecsdl.org/content/159/8/A1209.full )
and:
H2 + 2 OH- ⇒ 2 H2O + 2 e- (See Eq [2] above reference)

Cl- + H2O ⇒ HCl + OH-

At the Aluminum strip (cathode) we could have:

H2O2 + 2 e- ⇒ 2 OH-
Na+ + OH- ⇒ NaOH

As the NaOH and HCl recombine forming NaCl, its chief role here is to transfer charges, that is, acts as an electrolyte.

A significant side reaction forming Fe2O3.xH2O is:

2 Fe(OH)2 + H2O2 --> 2 Fe(OH)3

Now, as to whether this is an effective manner to prepare Fe2O3, I tested a reaction mixture of an piece of Iron in a NaCl + a dilute H2O2 solution in an aluminum pan. Within an a few hours, an obvious formation of a rusty water solution.

I also tested the reaction of Al pieces and a chunk of Iron in a solution of NaCl, NaOCl and vinegar (a so called Bleach battery if you substitutes Cu for Fe). Again, within a few hours an even more obvious formation of a fine rust suspension in water. This latter battery cell also forms, unfortunately, a significant amount of Chlorine.

Added references please see http://www.exo.net/~pauld/saltwater/ and also http://www.dtic.mil/dtic/tr/fulltext/u2/d019917.pdf

[Edited on 26-10-2013 by AJKOER]

AJKOER - 29-10-2013 at 05:55