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is there anyway to speed this reaction up at all?

 

I`ve one reaction going for 2 years now, and I had a look today and I finaly have FeCl2 crystals.

ok it was from when I was on holliday by the sea and tied a string of large powerfull magnets to a rock as the tide came in and collected the particles that stuck to it, they were washed and put in a sealed jar along with 30% HCl.

 

now I have 30% HCl and old iron screws in there, other than making them look a little cleaner, there`s no visible signs of a reaction, I added a few drops of Nitric acid to this (literaly 3 drops in 100ml of HCl) hoping something would "kick-off", Nada :(

 

I don`t want to warm it up as that`ll just waste the HCl as gas.

 

any ideas?

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This is a problem with dissolvingf many metals in acid. Although literature states for most transition metals that they dissolve in non-oxidizing acids, like HCl, only a few really do at an acceptable rate (Mn, Zn, Cd). Many other metals, such as Ni, Co, Cr, V do dissolve, but VERY slowly.

 

You could try and add some 30% H2O2 as oxidizer to your 30% HCl. This makes a potent mix, capable of oxidizing many metals at much higher speed. However, this mix oxidizes Fe to the +3 oxidation state straight away.

 

When mixing 30% HCl and 30% H2O2 use a volume ratio of HCl : H2O2 = 5 : 1. You do not have to fear for explosions as can occur with piranha solution, but be careful though. This mix is amazingly corrosive and it gives off some chlorine gas.

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I only have 9% H202 at 30 vols, and it has phenacitin and phosphoric acid in it, so I`de rather NOT contaminate it further.

 

of course I have to worry about explosions if they`re unpredictable (same as any Sane person), the Cl gas I can cope with. I have Red iron oxide here (Fe2O3) that I`de have liked to have used, but that`s useless in HCl and even Sulphuric.

 

I`m low on metal chlorides, so I figured since I have a gallon of HCl I may as well stock up again, I`ve a feeling this one`s going to take a while!

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woelen - doesn't heating the acid speed up the reaction? Also could you not powder the reactants to make them react faster? (Although this probably won't work in this case because the reactants are metals...)

 

Cheers,

 

Ryan Jones :)

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woelen - doesn't heating the acid speed up the reaction? Also could you not powder the reactants to make them react faster? (Although this probably won't work in this case because the reactants are metals...)

 

Cheers' date='

 

Ryan Jones :)[/quote']

Heating does speed up the reaction, but even then it still is quite slow. I did an experiment, similar to YT's, dissolving iron and titanium in conc. HCl. It really takes days before an appreciable amount of metal is dissolved.

 

H2SO4 is totally useless with these metals. The only reason, why they dissolve somewhat in HCl is that they form chloro-complexes, otherwise they would not dissolve at all.

 

In fact, I prefer powdered metals for my experiments, although I do not like the really fine powders, due to their pyrophoric nature. So, I have coarse titanium powder, coarse iron powder and coarse copper powder. I do not prepare stock amounts of metal chlorides, if I want to do an experiment with e.g. TiCl3, then I make it, let it react for a day or two and then I play around with it. This requires some planning ;) , but it also enforces you to think well in advance and generally improves the quality of the experiments :) .

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  • 2 weeks later...

A very odd reaction has occured, the setup as outlined in post #1 has changed color entirely, from the Bright yellow FeCl2 to a light Green clear soln and where`s a white PPT at the bottom?

there is clear evidence that the iron screws I put in have corroded a little also.

 

basicly dissolve some FeCl2 (obtained from sand in HCl) in a little water, add some HCl and then drop in an iron screw or nail, and then wait a week or so.

 

the only possible thing I can think of that May have done this, is that when I dissolved the magnetic particles from the sand, they weren`t all Iron and could be Nickel in there also (that would explain the light green color).

 

would anyone else agree with this hypothesis?

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A very odd reaction has occured' date=' the setup as outlined in post #1 has changed color entirely, from the Bright yellow FeCl2 to a light Green clear soln and where`s a white PPT at the bottom?

there is clear evidence that the iron screws I put in have corroded a little also.

 

basicly dissolve some FeCl2 (obtained from sand in HCl) in a little water, add some HCl and then drop in an iron screw or nail, and then wait a week or so.

 

the only possible thing I can think of that May have done this, is that when I dissolved the magnetic particles from the sand, they weren`t all Iron and could be Nickel in there also (that would explain the light green color).

 

would anyone else agree with this hypothesis?[/quote']

FeCl2 is not bright yellow, it is (almost) colorless. What you have is a complex, FeCl4(-), which is bright yellow.

 

In fact, all purely aquated iron ions, both [Fe(H2O)6](2+) and [Fe(H2O)6](3+) are almost colorless. Look at the solid salts of these. E.g. ferric nitrate or ferric sulfate are very pale, almost white, the same holds for ferrous salts. These are very pale green/blue.

 

I also made solutions of FeCl2 in conc. HCl with reagent grade iron powder and these solutions are very light green, almost colorless. In due time (weeks) they become yellow, due to slow absorption of oxygen from the air.

 

What you now observe is that the bright yellow chloro iron (III) complex oxidizes the iron metal, resulting in formation of iron (II):

 

2FeCl4(-)(yellow) + Fe ---> 3Fe(2+)(pale green) + 8Cl(-)

 

So, your bright yellow liquid becomes pale green.

 

If you remove the stopper, then the liquid becomes yellow again.

 

Have a look at my page on iron salts and look at all the colors. This page may help you to understand what you observed:

 

http://woelen.scheikunde.net/science/chem/solutions/fe.html

 

Take special notice of the colors of the aquated ions, the hydrolysed ions and the chloro-complex.

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Hey, sorry to resurrect this board, but it made more sense to add on here then make a new thread. I just decided that it'd be neat to make some Iron Chloride. I took steel wool (I don't have so many reagent grade chems) and disolved it in hydrochloric acid. Came back and it was yellow. Added more steel wool, it turned light green with white prec. Just like YT. I'm just trying to figure out how to actually get FeCl2 when my aolutions oxidize on me. Woelen said that the Oxidation takes weeks but within 4 hours the steel wool dissapeared and the solution was bright yellow. (Also @ YT: If you take care to clean off any oils put on the steel wool to keep it from oxidizing, it reacts very decently (meaning you can actually see hydrogen bubbles. At first nothing happens as a layer of oil,Fe2O3, or w/e is disolved)) So my question is, with my solution oxidizing all the time, how do I create FeCl2 instead of FeCl3?

 

EDIT: Also YT, did you say that you got FeCl2 from the Ocean? I beleive you made a typo but your message says that at some point. I think you got Fe3O4...

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How do you perform your reaction? Just with the vessel in direct contact with air. I did my experiments in a cork-stoppered test tube and then air only enters the test tube very slowly.

 

I'm afraid it is the same trouble as with CuCl and aqueous copper (I) compounds. FeCl4(-) is very stable, even at low concentration of chloride ion, so the iron (III) species is strongly favoured under your conditions. The situation is not as bad as with CuCl and the colorless CuCl2(-) solution, but still, it is annoying.

 

So, if you dissolve your iron in a solution with iron (III) and HCl and you close the vessel, and you use excess iron, then you can get iron (II) in solution and keep it like this. But every time, when you open up the vessel to take out some, it will become colored yellow again. Even small amounts of FeCl4(-) make the liquid appear yellow quickly.

 

You could delay the process by taking a can of 3M dust remover. This is a pressurized relatively inert gas (butane, dimethylether, or 1,1-difluoroethane), which is used for blowing dust from camera-lenses, computer keyboards, audiodevices, etc. Each time, before closing the vessel with the ferrous chloride solution you could blow in some of this gas and then very quickly capping the vessel. This strongly reduces the oxidation, but does not eliminate it.

 

I only know one iron (II) salt, which stores reasonably well and that is Mohr's salt, a double-salt, FeSO4.(NH4)2SO4.6H2O, ferrous ammonium sulfate. This salt is not oxidized by air. Plain ferrous sulfate, FeSO4.7H2O in due time becomes covered by a brown crust of basic ferric sulfate.

An unoxidized ferrous salt is very pale green/blue, see

 

http://woelen.scheikunde.net/science/chem/compounds/ferrous_amm_sulfate.html

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Yeah.. I know this much so far.. I found myself that the situation was very similar to my last one. my question was, how to crystalize the Iron Chloride? Also, you said that the precipitate I have in there was formed because of the Iron III correct? If so, then reducing the solution and decanting followed by your perging of oxygen in the container should get rid of the prec. right?

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EDIT: Also YT, did you say that you got FeCl2 from the Ocean? I beleive you made a typo but your message says that at some point. I think you got Fe3O4...

 

LOL, No, I got the Metal from the ocean by stringing a line of strong magnets around a large rock and waiting for the tide to come in, then took the particles off the magnets.

I then left these particles to dissolve in HCl for some time and got a bright yellow/orange clear liquid, like the ferric chloride used to etch PCBs with.

my question is based around Why it`s changed color when I added a couple of old (partly rusty) screws to the soln.

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WOW! I filtered my FeCl4- Solution jsutn ow. I'm left with a very clear dark yellow liquid. To my surprise, I saw that the precipitate I had seen included some crystals of FeCl2. At first I didnt understand how YT had FeCl2 in his solution of FeCl4- solution. I'm not positive how this happened though. Well. Maybe the solution was saturated in FeCl4-, But there was so much HCl left that it continued to Oxideze Iron, resulting in a precipitate of FeCl2. Could someone second this or tell me im wrong?

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WOW! I filtered my FeCl4- Solution jsutn ow. I'm left with a very clear dark yellow liquid. To my surprise, I saw that the precipitate I had seen included some crystals of FeCl2. At first I didnt understand how YT had FeCl2 in his solution of FeCl4- solution. I'm not positive how this happened though. Well. Maybe the solution was saturated in FeCl4-, But there was so much HCl left that it continued to Oxideze Iron, resulting in a precipitate of FeCl2. Could someone second this or tell me im wrong?

Ferrous chloride dissolves approximately at 10 gram per 100 ml of water, according to my data. This compound crystallizes as FeCl2.4H2O and in solution it forms ions Fe(2+) and Cl(-). 10 grams of this compound is approximately 0.05 mol.

 

This means that the solubility is approximately 0.05 mol per 100 ml, hence 0.5 mol/liter. For soluble compounds, there is the concept of solubility product. For FeCl2 this is (taking the value of 0.5 mol/l as the maximum amount which can be dissolved):

 

Ksp = [Fe(2+)]*[Cl(-)]*[Cl(-)] = 0.5*1.0*1.0 = 0.5 mol3/l3

 

Remember, for [Cl(-)] the double value must be plugged into this formula, because we have 2 mol of Cl(-) ions for each mol of FeCl2.4H2O

 

Now, the common ion effect comes into play. In HCl, there are LOTS of chloride ions. Suppose you start off with 30% HCl and half of its chloride remains, the rest is used up to form FeCl4(-). Then still 15% HCl remains. Taking into account the higher density of HCl-solution, 15% HCl roughly contains 170 grams of HCl per liter, which means roughly 5 mol/l.

 

Plugging into the formula for Ksp:

 

[Fe(2+)]*5*5 = 0.5 ==> [Fe(2+)] = 0.02 mol/l

 

Using the molecular weight for FeCl2.4H2O and the fact that one Fe(2+) ion results in one FeCl2.4H20, this means that only approximately 4 gram of this can be dissolved in one liter of 15% HCl.

 

I hope you grasp the concept of this computation. All the numbers I have taken are VERY rough estimates, but using this kind of reasoning you certainly can explain why you see solid FeCl2.4H2O. Due to the rough estimates I may be off by a factor of 2, but still then, the reasoning holds.

 

So, the answer is 'yes', you can certainly have solid FeCl2.4H2O in your liquid. What I think is more important, however, here, is that you understand this way of reasoning.

 

------------------------------------------------------

 

Just for education, a related little experiment you can do without problem (you have the reagents, I'm quite sure :) ).

 

  • Take 2 ml of water and dissolve some table salt in it. See how much you can dissolve.
  • Next, take 2 ml of concentrated hydrochloric acid and see how much table salt you can dissolve in that.

Try to explain the observation.

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The results of my experiment corrolated with what I found in the Iron Chloride solution. The NaCl disolved very Poorly in HCl as apposed to H20. So basicly, the HCl is saturated with Fe2+ and Cl- very rapidly, but solubility is aweful because of the amount of chloride ions already in solution as HCl. I noticed ungodly amounts of HCl vapor eminating from my reaction vessle. Much more than would be expected in fuming HCl by it's self. I hypothesize that when the point in time comes when the water is saturated in chloride ions and FeCl2 begins to precipitate, that HCl is also forced out of the solution proportionatly. This is of course an educated guess. The liberated Hydrogen could give me false judgment, but these fumes weren't clear like H2, they were white and smelled horribly of HCl. I'm going by the strenth of odor here... Anyhow, learned something ese new.

 

Next. I have no idea what you guys are talking about with your Iron not disolving at any appreciable rate. I figured steel wool wasnt tremendously pure, so i went and bought some of those steel scratch pads made from the coiled steel. Pop one of those in HCl and it bubbles very vigorously. (31% HCl).

 

In my last experiment, I added a little steel wool at a time and alowed it to Oxidize to FeCl4- each time. This time, I threw a large amount of iron (Which STILL all disolved) in HCl and let it react away. Some neat has happened. I see 3 distinct layers. One layer of dark green precipitate (FeCl2?). Next, a saturated solution of (FeCl2?), and on top, a lighter green precipitate (Hydroxide? something else?) The layer of precipitate on top of the water is NOT because of hydrogen sticking to it and causing it to float. I have shaken the mixture many times and within 10 seconds a break occurs between the two precipitates. Also, the neat thing is, this top layer of precipitate is capping the liquid and therefor stopping it form oxidizing. The residue on the bottle however is yellow from oxidation as there was nothing preotecting it. Could someone verify what this top precipitate is? I havn't isolated it yet, it could be white for all i know and the solution is staining it. If this was Fe(OH2), I would expect it to sink in water...

 

EDIT: Ok guys, I decided it was lazy of me to ask what this was when i had done no tests myself. So I set out to isolate the top prec. I put some in excess water in a test tube to wash out any FeCl2 on it. I shook it up and YES, a snow white precipitate floated up very rapidly. So, this compound is white. But what is it? In a whiel I will see if it was oxidized any, but I don't think it will, as its been sitting around a while already.

 

Also, I believe that Yt's iron was actualy Fe3O4, and that caused it to react slower. I've presented this before but he insists on "Metal" being deposited on the magnet, but if I remember correctly, Fe3O4 Is attracted to magnets also and I know for a fact it is abundant in beaches....

Just reread Yt's first post. I also had a hark time with Screws, but I did see bubbling. If the scres were coated in soemthing that may retard (heh) the reaction.

 

Uh oh.. I just re read the whoel thread and Woelen says Fe2+ is almost colorless, but I have a VERY dark green solution. I can't even get light through it very well.. Maybe my Iron was alloyed and the white Prec. as well as the dark green color was form a contaminent (I hope not.)

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Remember that the surface area of steel wool is a LOT greater than the surface area of an iron nail. Surface area plays a HUGE role in determining the rate of a reaction. (Hence why a platinum gauze is used to oxidize ammonia in the production of nitric acid and not just a lump of platinum metal. It's also why the palladium/rhodium/platinum in your catalytic converter is dispersed onto a honeycomb instead of existing as solid lumps of metal).

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well I can honestly say that these iron screws in 30% HCl did nothing at all in the way of visible reaction (at least non that you`de actualy sit there and see).

I`ve still got these screws in the soln and it has gone a little darker (but still clear) green, this has been over the time since the original post.

as I said, nothing you`de want to sit there and actualy Watch :)

no gas at all?

 

I`ve not tried the steel wool idea yet though (I`ll have to get some).

I`m glad you managed to get that dark yellow but clear Chloride though, at least I know it wasn`t something I`de done wrong, not there`s much anyone CAN do wrong with a 2 ingredient experiment :)

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not there`s much anyone CAN do wrong with a 2 ingredient experiment :)

 

ha.. Yeah I replicated your dark yellow solution with the steel wool but the scratch pad thing is being weird. I set some of THAT green solution out and it's yet to turn yellow. Oh yeah.. Let me add some HCl... lol Well, We'll see what happens but I still dont know what the white floaty prec. is! I dried some out and it dried off white. It disolves in NaOH but not HCl, but looking at some solubility rules; There isnt a compound that wont disolve in HCL but will in NaOH except I believe Thallium. This leads me to believe that Adding this white compound to NaOH does not change the composition of the mystery compound, rather that it only disolves under alkaline conditions. Either that or not ;), heh. Also I might add, adding HCl tothe disolved material does not yeild a precipitate. Therefor the beginning white compound was not a chloride.. What other anion could it be?

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ha.. Yeah I replicated your dark yellow solution with the steel wool but the scratch pad thing is being weird. I set some of THAT green solution out and it's yet to turn yellow. Oh yeah.. Let me add some HCl... lol Well, We'll see what happens but I still dont know what the white floaty prec. is! I dried some out and it dried off white. It disolves in NaOH but not HCl, but looking at some solubility rules; There isnt a compound that wont disolve in HCL but will in NaOH except I believe Thallium. This leads me to believe that Adding this white compound to NaOH does not change the composition of the mystery compound, rather that it only disolves under alkaline conditions. Either that or not ;), heh. Also I might add, adding HCl tothe disolved material does not yeild a precipitate. Therefor the beginning white compound was not a chloride.. What other anion could it be?

You made me curious with all your iron experiments and I did an experiment with reagent grade HCl, 30% by weight and 99.9+ % pure iron powder, average particle size appr. 100 μm.

 

I did a small spatula full of iron powder in a test tube and added some of the HCl. As soon as I did this, there was some slow bubbling, but definitely not vigorous. The liquid almost at once turned light yellow. On standing for a while, the yellow color turned somewhat more intense. Having the mix stand like this would take ages before all of the metal is dissolved.

 

Next, I heated the test tube to appr. 90 C (not boiling, well below that point). At this temperature, the bubbling became much faster, but still not very violent.

While the liquid was bubbling nicely, I capped the test tube loosely, in order to allow the gas to escape, but not letting air in. I heated for several minutes and it took me more than 10 minutes to dissolve the small amount of iron in the acid. At the point, where it was almost gone, I heated more strongly, until the liquid boiled vigorously and then at once, I took it out of the flame and capped it tightly (it still was loosely capped). In this way, I assured that no air is left in the test tube and on cooling down there is under-pressure.

 

The liquid now almost is very pale green/blue, with a fairly large percentage of iron dissolved. All iron is dissolved as iron (II) species. On cooling down, the liquid remains clear. No flocculent precipitate can be observed.

 

My observations can be explained as follows:

 

First, in the HCl some oxygen is dissolved. This causes formation of the yellow FeCl4(-) complex (iron in oxidation state +3). When more iron dissolves and there is no air allowed to enter the test tube, air being displaced by hydrogen gas, then the iron (III) species is reduced to iron (II) species again. On longer standing, nothing changes anymore. So, the capped test tube contains a light green/blue liquid, which is totally clear.

 

Some small pieces of metal remain undissolved. At a certain point the activity of the liquid (the acid) becomes so low, that the metal does not dissolve anymore.

 

I attached a picture of my final result. It nicely shows the green solution.

 

I could not reproduce your results of dark green liquids, flocculent precipitates etc. So, I'm afraid that your source of iron is quite impure. If it contains chromium, then a green color is understandable.

ironII.jpg

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Alright.... Wel thanks. Your results differ vastly from mine. I will look for a more pure source of Iron.... AS a side note. I have a large quantity of CuCl2 solution evaporating and I see some beautiful crystals. I could never get the CuCl2 pure, it was always very green. I found out that the reason it was so impure was I was boiling it. I let my CuCl crystalize over the course of a few weeks and its hardly green at all. It looks about the same color as your reagent grade. Like an aquamarine color maybe... well bluer than that. Just an update...

 

 

also. I told you guys no precipitate was formed when putting HCl in the NaOH mystery liquid solution.. well i cam home today to find white crystals... Some are cubic and some are stick like... lol I may never know what that is...

 

What WOULD be a pure fom of iron I could buy without spending lots.. Is steel wool decent enough it it is washed thoroughly?

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Getting really pure iron at some local store is not that easy. The reason for that is that very pure iron hardly has any commercial application. It is not sufficiently hard and strong for construction purposes and it rusts away easily. So, commercial iron is alloyed with all kinds of elements in order to prevent rusting and to make it stronger (e.g. chromium, vanadium, nickel, carbon). Steel wool, as the name says, is steel, which contains a LOT of other elements besides the iron (mostly carbon, but also other materials).

 

From online shops, you can obtain iron at very high purity easily, but that requires you to buy online and pay with PayPal or credit card. The price should not be prohibitive

 

http://www.emovendo.net/store/customer/home.php?cat=253

 

Iron is present over there at just $6 for 30 grams of ultrapure metal (99.98+ % purity). From this company I have some of the cheaper ultrapure elements (Fe, Se, Ag, Pb, Al, Sb). Shipping is only $1.60 to me in the Netherlands, so for you it should be even cheaper.

 

Sometimes, eBay also has pure iron. My powdered iron is from eBay, at a price of EUR 7 per 100 gram, but purity is only 99.9+ % instead of 99.98+ %. That 0.1% can make quite some difference, especially if the impurity is a strongly colored species. Fortunately for my iron this is not the case, the impurity is mostly carbon, making the iron a little less reactive and causing the final pieces of metal not to dissolve as my picture shows (I assume they have a high carbon content).

 

If you want to do nice metal experiments, then I also would suggest you to do experiments with chromium. That element gives beautiful blue and green solutions. Pure chromium also can be had from eBay sometimes. Emovendo, unfortunately is quite expensive on chromium, but with some luck, you can find cheaper sources of very pure chromium.

 

Can you make a picture of your nice blue CuCl2.2H2O crystals? It is interesting to read that you get very pure CuCl2, which is not green. Indeed, slow evaporation of solutions is a good way to get nice crystals, the only disadvantage being that it may take VERY long. Most people are not that patient :D .

 

The cubic crystals you get most likely are NaCl (from NaOH + HCl). Needle-like crystals? I cannot answer that, because I do not know the impurities from your iron and solutions.

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well here`s an unexpected surprise, my iron in HCl has now developed a distinct orange band just above the liquid level, strongly resembing Rust, the liquid itself is still a green color, but has some orange tinge in it now also.

I think it`s perhaps gone Too far and started breaking down somehow?

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Good Call on the Crystals, that'd make sense. I still have some liquid left in my tub of CuCl2. I know that it becomes pure because I set aside smaller amounts to crystalize and they showed promise. It may be another week before my large tub evaporates, but I will let you know when this happens.

I threw a screw that I had in acid already prevoiusly to get rid of any coating. After Putting in a test tube with HCl, The Fe2O3 layer melted off (this was actually kind of neat) and the solution turned yellow. All at once it bubbled decently quickly and eventually the solution turned clear and then light green. I left it to react over night and even though the reaction vessel was loosely capped, the solution is now back to yellow. I also have a decent amount of whitish precipitate on the bottom, but no floaty crap! This seems to be a lot more pure than the scrubby stuff. So it looks like the white stuff is my FeCl2.

 

As a side note, it amazes me at the difference in composition of two "iron"s It would make sense, because the scrubby iron needs to be very bendable without breaking. Iron itself would not fulfill this purpose. So the dark green liquid, white floaty precipitate are apparently added metals to increase tensile strength.

 

Also, is my precipitate of FeCl2 decently pure? You talk about Fe3+ hydrolysing on your website. Your website gives me the impression that FeCl2 is decently Stable under neutral conditions. If I washed the precipitate in a liquid insoluble by FeCl2 and then added distilled water to disolve the FeCl2, colul I then crystalize it leaving the other crap behind? Or would this ALSO oxidize.

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