Out of curiosity I wanted to see whether I could establish how far along the oxidation process the FeCl3/Cl2 solution is to completion.
I decided to pipette unknown molarity of Sodium Carbonate into a sample of my Fe2+/Fe3+ solution.
Upon dripping Sodium carbonate solution into test tube: the first reaction was a very clear lime green/yellowy fluid suspended on top of the bright
red/brown/yellowish original solution which was being pushed down. As more Sodium carbonate was added a dirty olive green granular precipitate formed
on top of the clear lime green/yellowy fluid and I ended up with three suspensions in the test tube.
My understanding is that the dirty olive green is Fe2+
The middle is Fe3+
the bottom is my original solution
It only remains for me to ask if there is a method to take this one step further and determine actual ratios, how much Fe2+ is in the solution?
I just went outside, came back in after two minutes and the top layer just turned orange. Interesting but Not to sure why?
[Edited on 25-4-2014 by CHRIS25]blogfast25 - 25-4-2014 at 07:25
Fe2(CO3)3 in all likelihood doesn't exist in STP, watery conditions (see also similar non-existent carbonates of other trivalent ions like Al(III) and
Cr(III)). Instead Fe(OH)3 forms, and CO2.
Measuring the Fe(II)/Fe(III) ratio is seriously difficult, I think. Fe(III) could be determined by titrometry. Then determine total Fe = Fe(II) +
Fe(III). Algebra then gives you Fe(II)/Fe(III).
Potentiometry might also do it but would require standard solutions with known Fe(II)/Fe(III) ratios for a calibration line.
Oxidation of Fe(II) with peroxide in acid conditions is very fast, almost instantaneous. So assuming you use the right amount of peroxide you can
reasonably assume that at the end of this addition all Fe is Fe(III).CHRIS25 - 25-4-2014 at 09:15
Understood. It is just that I read that NaCO3 will cause a green precipitate if there are Fe2 ions and a light brown indicates Fe3. Since I got two
very clear distinct suspensions that this was indicating the two Fe Types. Although the middle layer was more yellow. I was trying to understand the
separation. Vigorous shaking turned the whole solution exactly the same colour as the middle layer in this image.jwpa17 - 25-4-2014 at 18:45
I think it's possible to determine the ratio of Fe(II) to Fe(III) using potentiometry - measuring the voltage of that redox couple versus a reference
cell. I'm sorry, but I don't have access to the article, but this seems to be appropriate: http://www.sciencedirect.com/science/article/pii/S1572665710...
and perhaps the attached paper. Good luck.
Over my head, the paper would make for an interesting read - but I won't buy it. The most advanced machine I have in my lab is a refrigus machine for
crystallizing dihydrogen oxide. I was interested in the three or rather two suspension separations really, are they what I said? and any ideas over
the orange and then the final mix I believe would have been the colour that indicates Fe3+ oxidation brought on by the NaCO3? I have still not
delved deeply into predicting reactions via understanding ionic and covalent equations. So maybe if I understand these better I could say what was
going on in the test tube. bismuthate - 26-4-2014 at 03:58
Well I believe that FeCl3+3 KI--->FeI2+3 KCl+1/2 I2 or FeCl3+KI--->FeCl2+KCl+1/2 I2
so you could use an iodometric titration if you wished.
Watch of for that evil dihydrogen monoxide .CHRIS25 - 26-4-2014 at 04:25
"Watch of for that evil dihydrogen monoxide" - yes, it's quite toxic when mixed with alchohol!
Sounds excellent, KI, now to get my hands on some of it. What is "12" in the above? was that supposed to be O2bismuthate - 26-4-2014 at 04:28
Its iodine. I2 not 12.CHRIS25 - 26-4-2014 at 04:42
Ah, ignorance is my greatest gift.CHRIS25 - 26-4-2014 at 04:59
Since never having done this particular titration before I needed to research, I came across this: http://www.metrohm.com/com/Applications/methods.html?identif...
Interested to hear your perspective. Obviously it is thoroughly accurate due to its source, but still, I wondered what you might say seeing this
would be the third time I found that titrating for Fe3+ ions in he presence of Fe2+ ions requires sodium thiosulphate as well. (I have both by the
way so I will get to work).bismuthate - 26-4-2014 at 05:15
Well it does require thiosulfate because you use that to react with the iodine and then you measure how much it took to neutralize it which will tell
you how much iodine and in turn Fe3 there is.CHRIS25 - 26-4-2014 at 05:43
Yes I see that, but what is this starch as an indicator, I know what starch is, but I have never heard of a special starch indicator. I see that corn
or potatoe starch can be used. Just any old starch?
[Edited on 26-4-2014 by CHRIS25]bismuthate - 26-4-2014 at 06:14
Just regular starch it forms a very deep blue/purple complex with iodine you keep and adding thiosulfate solution until the color goes away.blogfast25 - 26-4-2014 at 07:40
C25:
Yes, good find. Metrohm is a very good source of titration info. A while back, on request they sent me two handy booklets/manuals on just about any
titration imaginable.
This is a typical iodometric titration, of the type I was alluding to in my post. Ferric ions oxidise iodide to iodine, the iodine is titrated with
thiosulphate, using starch as indicator (it forms an almost black complex with iodine).
Ferrous ions don't interfere, in fact during the first step ferrous ions are formed.
The method could also be used to determine total Fe, by first oxidising the ferrous to ferric ions. Then titrate to get total Fe.CHRIS25 - 26-4-2014 at 10:38
Thanks Blogfast. Quickie, I read that Starch solution will decompose after a few days, but can add salicylic acid as a preservative, is this
something you are aware of?
[Edited on 26-4-2014 by CHRIS25]blogfast25 - 26-4-2014 at 12:45
Thanks Blogfast. Quickie, I read that Starch solution will decompose after a few days, but can add salicylic acid as a preservative, is this
something you are aware of?
[Edited on 26-4-2014 by CHRIS25]
There's all kind of 'stabilisers' for starch. I use pre-stabilised for my work. If you only need a bit, just follow an internet recipe to prepare a
bit from potatoes: it'll stay ok for at least a week (after that it'll mould or ferment).
.