June 10, 2005

    Today was a few assorted tries at this chemical.  I had some success at a previous time by dripping ferric nitrate into basified NaOCl, this formed colloidal Fe(OH)3 which was quick to react with the NaOCl (~10% Pool Chlorine).  The mixture was heated to 100C for a few moments and viola, Na2FeO4 which I precipitated with BaCl2.  However today I decided to try a modification to this.  I took 200 ml 10% NaOCl and added to this 15 g KOH and 5 ml Br2 to make the solution more concentrated in terms of hypohalite (considering the normal called for amount is a 15% solution).  FeCl3 solution (simple OTC circuit etching) was added drop wise with stirring until 15 ml were added.  Stronger magnetic stirring should have been necessary.  The solid wasn't quite colloidal but I heated it to ~75C and held it there for a few minutes, the solid turned black.  This was all inspired due to a recent paper I read stating that ferric nitrate was superior for this reaction (a fact I was aware of) but that it was superior because the oxidation occurred at a lower temperature (which I didn't know, the oxidation of FeCl3 supposedly doesn't happen appreciably with NaOCl until about 70C).  So, the solution was filtered through glass wool and most of the colloidal Fe(OH)3 came through, I allowed it to settle and the solution was red, it was not the purple-red color I associate with the ferrate anion.  So I scrapped it.  If it would have worked then I would have filtered again and added a concentrated KOH solution, in the presence of this the solubility of K2FeO4 goes down and it precipitates but after a certain concentration of KOH is added then KCl will precipitate too.

    Anyway, after that failure I decided to go for broke.  I took 100 ml distilled H2O and added 25 g KOH to it.  The solution heated somewhat significantly and 3.5 grams Fe2O3 was added.  Most of this stayed in the water area but some settled.  To this was added drop wise 24 ml Br2 over the course of several minutes, initially on the addition of the first few drops a crackling sound would become apparent but it became less so as more Br2 was added.  The solution was heated upon finishing and filtered, nothing that resembled ferrate was obtained.

    A few days ago I made some nickel oxalate by dissolving 20 grams of oxalic acid dihydrate in 125 ml of boiling water, this was added to a solution prepared by dissolving 18 grams of nickel chloride (anhydrous) in 100 ml H2O.  A precipitate formed soon thereafter and this was allowed to settle for nearly a week until today.  It was filtered through two coffee filters under vacuum but about 25% of the solid made it through, the rest was a somewhat sticky thick mass that I added to a beaker to hold onto.  The stuff that made it through was set to the side to age and await filtration again at a later date.  

    I attempted to precipitate ferric oxalate by a similar procedure mixing ferric chloride and oxalic acid.  However very little precipitate was made, there are a number of reasons for this, the solution of ferric chloride itself was initially acidic due to the nature of FeCl3 this works against the reaction.  Additionally oxalic acid working with the Fe3+ cation can do some other things besides forming an insoluble compound, the oxalate can simply bridge to Fe3+ cations or can just act as ligands that do not cause the precipitation of the compound.  Never the less, the crystals that I do see on the bottom are brown/red and translucent so the color may just be due to the dark red solution.

June 11, 2005

    In an iron crucible (pictured in the next experiment) 10 grams of potassium hydroxide and 7 grams of sodium nitrate were combined with 2 grams of iron oxide.  The mixture was heated till it was molten using a propane torch and held there.  If the temperature was but slightly higher then the melting point the liquid would bubble and froth and go over the edges, a larger container is definitely necessary.  The mass started red but eventually settled on the brown color in the picture above.  Upon dissolving in distilled water no coloring is evident in the water.  So I would conclude little to no ferrate was formed.  I believe this is due to my inability to use higher temperatures then the melting point due to fear of it boiling over.  So I will at a future date take a piece of the above and heat to a higher temperature hoping for the characteristic color of ferrate.

    Ten grams of sodium chloride and 10 grams of calcium chloride were placed into a crucible and they were moistened with maybe 2 milliliters of water.  The mixture was heated with a propane torch and the entirety of it became liquid in a short time.  There was some spattering but for the most part it appeared homogenous when it was all over.  What I had hoped for is a quick transition from the liquid to the solid and back to the liquid phase followed by some experiments with electrolysis (note the anode attached to the top prong of the crucible at the top of the picture).  However it would not melt, no matter how long it sat over the blasting flame.  The eutectic mixture wasn't exactly where it should have been to be perfect so that's part of it and there was no insulation.  Even later when I placed a clay tile over the top it wouldn't melt.  It seems getting this mixture to melt is harder then I had anticipated.

June 12, 2005

        Going for simplicity the reaction mixture at left was produced by adding 5 ml concentrated ferric chloride solution to the beaker and immediately adding 150 ml 10% NaOCl to it and swirling.  The mixture was heated on a hotplate to boiling when choking fumes came off it then I sat it to the side to settle.  Meanwhile the second beaker was produced by mixing 125 ml NaOCl with 10 grams NaOH and allowing to dissolve (during which time much heating took place) and thereupon 5 ml concentrated ferric chloride was added.  This was heated a short while on a hot plate and when it was done settling it looked about the same as it does in this picture.  The solution on the left that was produced in the easier manner of the two actually looked to contain some ferrate, when filtered (see picture to the right.) I was left with a purple solution.  The other solution on filtering gave a clear solution.  More experiments are necessary.

 

 

    Using the nickel oxalate that I made on June 10th (see above), I attempted to make nickel powder today.  I took the oxalate and placed it in a jar that I had for the storage of reagents and started heating.  A large quantity of water came off and the pieces clumped together, I found that a test tube broke up the hot blue pieces easily so I continually crushed and stirred them as they heated.  I came to the conclusion it would be better to simply heat the nickel oxalate on a boiling water bath or oil bath then crush it in a mortar and pedestal before attempting to pyrolyze it but that conclusion came too late to be of much help.  Eventually though I found to my dismay that the container I was using to decompose the oxalate was no Pyrex, Kimax, or any other variation that proved to have heat resistance as it cracked loudly up the side and threatened to fall apart right when the heat started to really decompose the oxalate (note, when grandules of it fell on the hotplate the decomposed and caught fire leaving a black dust in their wake, I would have hated for the whole mass (>20 grams) to meet the same fate.  So I removed the container and transferred it to a Kimax volumetric flask and loosely stoppered it.  The powder continued to give off water and as it did so I continued to swirl the powder which resulted in it clinging to the walls.  Eventually most of the powder at the bottom was decomposed but there was still adherent unreacted powder on the sides of the flask, I used a propane torch and fanned it over the walls to decompose that oxalate as well.  Eventually I was left with a volumetric flask full of black powder (see picture below), now to react it with aluminum powder in an inert atmosphere.

Current Comments

2 comments so far (post your own)

how can one make FeCl3?
Also, what are the main uses of ferrates?

Posted by b on Tuesday, 06.27.06 @ 19:07pm | #31

To make FeCl3 you could just dissolve iron in some hydrochloric acid and let it sit with exposure to air. The oxygen is a strong enough oxidant to bring the Fe2+ to Fe3+ providing there is excess acid. This will make the hexahydrate if it is carefully heated to drive off water or better yet if a vacuum is used. Over heating though can result in decomposition of the hexahydrate to give HCl among other things.

Ferrates are hardly used in chemistry, they don't have very good shelf lifes and are fairly strong. I have seen their use in waste water treatment though.

Posted by BromicAcid on Tuesday, 08.1.06 @ 19:44pm | #34

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