Sciencemadness Discussion Board - Test for Fe2+ and Fe3+

Test for Fe2+ and Fe3+

agent_entropy - 16-6-2015 at 12:29

I'm trying to put together a qualitative test kit for iron in water leached from soil samples that can detect both Fe2+ and Fe3+. My idea is to have a bottle of dilute nitric to acidify the sample, then a solution of potassium thiocyanate to test for Fe3+. If that fails to yield the red color of a positive test (ferric thiocyanate complex), then add sodium chlorate to oxidize any Fe2+ to Fe3+ and produce the red color.

The test seems to work well enough, but my question is, does anyone see any danger in that? Potential for formation of dangerous (explosive) chlorate salts in the waste disposal?

I realize hydrogen peroxide would serve adequately as oxidizer, but I'd like to avoid samples foaming. I picked chlorate because of availability and the fact that it's colorless (also because nitric failed to oxidize Fe2+ in dilute solution).

kecskesajt - 17-6-2015 at 00:45

I dont think so.Soil dont have so muck reducable components to make any explosive product.And in water, nitrogen triiodide is stable so keep wet the samples.

woelen - 17-6-2015 at 04:24

You can use a solution of potassium hexacyanoferrate(II), a.k.a. yellow prussiate of potash for detecting the iron. This test is VERY sensitive, more so than the thiocyanate test. It also requires very small amounts of reagents.

In practice the test works as follows:
- Add your sample to 100 ml of water.
- Add a few drops of dilute acid (e.g. 10% HCl or 10% HNO3) to your sample
- Add a drop of 3% H2O2 to 100 ml to your sample, stir and allow to settle for a minute or so
- Add a few drops of a 10% solution of yellow prussiate of potash.

If any iron is present (either as iron(II) or as iron(III)), then the liquid turns bright blue.

The reagents for this are easy to obtain. The yellow prussiate of potash stores well. No need to worry about foaming or strong acids, because the acid reagent and hydrogen peroxide are used at great dilutions.

agent_entropy - 17-6-2015 at 08:17

Thanks woelen! I knew there was a test for iron using hexacyanoferrate but I was under the impression that it only detected iron(III). If it detects both, then I'll go with that.

papaya - 17-6-2015 at 08:44

Btw, woelen, why hexacianoferrate (II) detects both irons? what about hexacyanoferrate (III) ? (at least I know the latter detects Fe (II) )

Ozone - 17-6-2015 at 19:21

I used to do this with o-phenanthroline. An aliquot is treated with o-phenanthroline to make [Fe(phen)3]2+ (ferroin). The absorbance is measured with a spectrophotomer vs. external standard curve (I used ferrous sulfate). This gives Fe2+.

Another aliquot is treated with hydroxylamine to reduce any Fe3+ to Fe2+, treated with o-phenanthroline, and the aborbance is measured. The total iron is thus quantified, and Fe3+ can be given by difference.

See, for example: http://www.ebay.com/itm/1-10-Phenanthroline-monohydrate-o-Ph...

O3

[Edited on 18-6-2015 by Ozone]

woelen - 17-6-2015 at 23:11

Quote: Originally posted by papaya

Btw, woelen, why hexacianoferrate (II) detects both irons? what about hexacyanoferrate (III) ? (at least I know the latter detects Fe (II) )

Hexacyanoferrate(II) detects iron(III), but the drop of H2O2 (3%) oxidizes any iron(II) to iron(III), so this makes the test sensitive for both iron(II) and iron(III).

The hexacyanoferrate(III) (the red salt) detects iron(II), so, using a reductor, together with the red prussiate would be an option as well. The use of the yellow salt, however, is preferrable. The red salt is light sensitive and solutions of it slowly decompose, giving many kinds of complicated dark blue cyano-complexes (no free cyanide is produced), while solutions of the yellow salt are more stable and the solid keeps indefinitely, both in the dark and under light.

agent_entropy - 18-6-2015 at 08:56

woelen, you mentioned that hexacyanoferrate(II) is more sensitive than the thiocyanate test, any idea what the detection limit is? ie. How low can the iron content be before the blue color is either imperceptible or could be easily missed?

Ozone, thanks for the explanation of the o-phenanthroline test, I have a version of that from Hach. I was just looking for a quicker/simpler qualitative version.

aga - 18-6-2015 at 12:04

Having been handed the solution to your question on a plate, a little experimental work on your part would quickly give you the answer to your own question no ?

Please do those experiments, document them, and post the results here.

I for one would be very happy to have all that careful work done for me so i would not have to expend any effort at all.

Edit:

If you are unsure of the process, discuss it here for advice, and i will spend the time to replicate the (eventual) process to confirm your results.

[Edited on 18-6-2015 by aga]

woelen - 19-6-2015 at 01:33

Quote: Originally posted by agent_entropy

woelen, you mentioned that hexacyanoferrate(II) is more sensitive than the thiocyanate test, any idea what the detection limit is? ie. How low can the iron content be before the blue color is either imperceptible or could be easily missed?[...]

I do not know the exact figures, but one important difference between the two tests is that the thiocyanate test requires fairly high concentration of thiocyanate. The complex with thiocyanate is labile and when the concentration of thiocyanate is low, then it tends to decompose again in its separate ions. The hexacyanoferrate complex, on the other hand, is very stable and once it is formed, it does not decompose anymore, provided the pH is kept on the low side of neutrality (hence the need to add a few drops of acid).

In practice, I have done this test myself and it is amazingly sensitive. If you take one little crystal (appr. 1 mm diameter) of an iron salt and dissolve that in a liter of water and then take 100 ml of this, add a few drops of acid, a drop of hydrogen peroxide and a small spatula of hexacyanoferrate(II), then you get a clearly visible blue color.

It should be possible to use colorimetric determination for more quantitative results, but for a simple test, which is performed on location, this of course is not possible, when you just have a small test kit with you.

blogfast25 - 19-6-2015 at 02:37

Quote: Originally posted by woelen

The complex with thiocyanate is labile and when the concentration of thiocyanate is low, then it tends to decompose again in its separate ions.

I don't think that is strictly speaking true.

The FeSCN2+ has two disadvantages compared to Prussian Blue.

1. The formation constant (Kf

of FeSCN2+ is quite low, in the order of 400 to 500, IIRW. Prussian Blue is insoluble, so it has a much higher Kf.

2. Prussian Blue is very intensely coloured. much more intense than FeSCN2+(#).

Both points makes Prussian Blue much more sensitive. And of course it allows detection of both Fe(+2) and/or Fe(+3), assuming you have both versions of
hexacyanoferrate.

(#)The full structure is likely to be [Fe(SCN)(H2O)5]2+, I think...

[Edited on 19-6-2015 by blogfast25]

Augustus - 23-6-2015 at 12:57

Such a test will not work with either thiocyanate or hexacyanoferrate(II). Since water was in contact with soil and it's pH was 5 or higher then almost all iron was absorbed by soil as Fe2O3*nH2O (don't even hope to find iron(II) in there). Concentration of iron in this water will never be enough to be detected by thiocyanate or hexacyanoferrate(II). The only solution is to leech iron from soil samples with acid.

aga - 23-6-2015 at 13:04

Are you speaking from experience ?

If so, please expand on how/what you did in your soil Fe measurements/xperiments.

If not, erm, what ?

Augustus - 23-6-2015 at 13:18

Yep, from experience.
Second solution is to use solvent extraction (100 ml water and 1 ml CHCl3) and a reagent of choice for iron.

Augustus - 23-6-2015 at 13:22

I used to detect heavy metals with ppb dilution in drinking water with dithizone and CHCl3 extraction.

aga - 23-6-2015 at 13:27

If you could elaborate, and explain how you did those procedures, i'm sure that would be of great help to the OP, and of interest to all of us.

[Edited on 23-6-2015 by aga]

blogfast25 - 23-6-2015 at 14:03

Quote: Originally posted by Augustus

Such a test will not work with either thiocyanate or hexacyanoferrate(II). Since water was in contact with soil and it's pH was 5 or higher then almost all iron was absorbed by soil as Fe2O3*nH2O (don't even hope to find iron(II) in there). Concentration of iron in this water will never be enough to be detected by thiocyanate or hexacyanoferrate(II). The only solution is to leech iron from soil samples with acid.

He would acidify the soil sample. That should be enough to leach out Fe(+3).

aga - 23-6-2015 at 14:06

Quote: Originally posted by agent_entropy

a qualitative test kit for iron in water leached from soil samples that can detect both Fe2+ and Fe3+. My idea is to have a bottle of dilute nitric to acidify the sample

I guess the OP has that aspect covered.

It would still be great to hear how ppb detection works.

blogfast25 - 23-6-2015 at 15:44

Quote: Originally posted by aga

It would still be great to hear how ppb detection works.

Not putting my hand in the plasma on that but ICP atomic emission spectroscopy might do it.

Having said that it would be interesting to see how low Fe(3+)/K4Fe(CN)6 goes. Not hard to do. Might be able to find it in a nook or cranny of the Tinkernettings...

[Edited on 23-6-2015 by blogfast25]

agent_entropy - 24-6-2015 at 04:49

I'm actually hoping to NOT find free iron in the soil leachate (cleaning up an iron (III) chloride spill) so if it is indeed absorbed as Fe2O3*nH2O that would be wonderful. However, I still need to do the tests in order to quell the concerns of others.

@aga, my potassium hexacyanoferrate (II) finally arrived today and I will begin experimenting both visually and spectrophotometrically to try to find a detection limit.

aga - 24-6-2015 at 04:57

Cool.
photos would be nice along with your findings.

agent_entropy - 24-6-2015 at 06:11

I just realized that spectrophotometric investigation will be essentially meaningless since Prussian blue is insoluble. I guess I'll continue visually.

[Edited on 24-6-2015 by agent_entropy]

blogfast25 - 24-6-2015 at 07:17

Quote: Originally posted by agent_entropy

I just realized that spectrophotometric investigation will be essentially meaningless since Prussian blue is insoluble. I guess I'll continue visually.

[Edited on 24-6-2015 by agent_entropy]

At low concentration it is colloidal. It will still colour a solution, even at very low concentrations. There are colorimetric determination protocols based on Prussian Blue, if I'm not mistaken.

agent_entropy - 24-6-2015 at 09:56

Indeed, Prussian Blue seems to be soluble or at least colloidal at low concentration.

Dropper bottles of testing reagents were prepared. 3M Nitric acid, 3% Hydrogen peroxide, 0.2M potassium hexacyanoferrate (II) (10% as the trihydrate), and 0.2M KSCN.

A 500mL stock solution containing 10ppm iron as Fe3+ was prepared volumetrically by dissolving 0.113g Fe(NO3)3*9H2O. The stock solution was diluted to 3, 2, 1, 0.5, 0.25, and 0.1 ppm standards.

5mL of each standard was pipetted into a small vial.

2 drops of 3M Nitric acid were added to each vial.

2 drops of 3% Hydrogen peroxide were added to each vial.

1 drop of 0.2M potassium hexacyanoferrate (II) was added to each vial.

About 10 minutes later the 10ppm sample had a precipitate of Prussian Blue.

The same procedure was repeated except that 1 drop of 0.2M potassium thiocyanate was added to each vial.

Visually, I would say that the hexacyanoferrate test becomes unreliable at less than 1ppm and the thiocyanate test becomes unreliable at less than 10ppm.

agent_entropy - 24-6-2015 at 10:00

UV-Visible spectra were recorded for both tests.

Hexacyanoferrate Test:

Thiocyanate Test:

It was also possible to construct a calibration curve for each test reagent.

Hexacyanoferrate Curve (using the 700nm absorbance):

Thiocyanate Curve (using the 430nm absorbance):

It would seem that, obviously, the spectrometer can determine colors better than the eye.

Also, no iron was detectable in any of the soil sample leachates.

[Edited on 24-6-2015 by agent_entropy]

woelen - 24-6-2015 at 22:59

Thumbs up, very nice to see such quantitative results!

What I find remarkable is the strong yellow/brown "background" color of all samples to which hexacyanoferrate(II) was added. This looks like oxidation to hexacyanoferrate(III). You said you added two drops of H2O2 (3%) to each sample and one drop of solution of hexacyanoferrate(II). Maybe you had too much peroxide? This test requires both iron(II) and iron(III). The iron(II) must be in the hexacyanoferrate. If all hexacyanoferrate is oxidized to the +3 state, then no blue complex is formed anymore, but a brown FeFe(CN)6 complex. This complex is fairly dark brown, but it certainly is not as strongly colored as the blue FeFe(CN)6(-) complex.

You could try adding just a single drop of H2O2 to each sample, and adding 2 or 3 drops of cyanoferrate(II) to each sample, just to be sure that some iron(II) remains. Maybe you can increase the sensitivity somewhat.

[Edited on 25-6-15 by woelen]

aga - 24-6-2015 at 23:11

Fabulous work agent_entropy !

It is very encouraging for us less experienced people to see how things Should be done.

Great images too.

blogfast25 - 25-6-2015 at 05:50

Yes, nice work indeed. Nice to see something quantitative here for once.

agent_entropy - 29-6-2015 at 05:30

I had to be out of the lab for awhile while my first son was born.

<-- shameless bragging

Anyhow, after sitting undisturbed for 111 hours the samples looked quite different. There was a blue precipitate in all of the vials subjected to the hexacyanoferrate test and there was no color in any of the vials subjected to the thiocyanate test.

2 drops of 0.2M KSCN were added to each of the thiocyanate vials, and the red color returned immediately.

[Edited on 29-6-2015 by agent_entropy]

agent_entropy - 29-6-2015 at 05:39

Per woelen's suggestion:

A new set of samples was prepared, again with 5 mL of standard solution each, at 3, 2, 1, 0.5, 0.25, and 0.1 ppm Fe3+.

One drop of 3M Nitric acid was added to each vial and swirled to mix.

One drop of 3% Hydrogen peroxide was added to each vial and swirled to mix.

2 drops of 0.2M Potassium hexacyanoferrate (II) were added to each vial and swirled to mix.

One additional drop of 0.2M Potassium hexacyanoferrate (II) was added to each vial and swirled to mix.

Another additional drop of 0.2M Potassium hexacyanoferrate (II) (for a total of 4 drops) was added to each vial and swirled to mix.

The same yellow coloration as before developed in each vial.

[Edited on 29-6-2015 by agent_entropy]

agent_entropy - 29-6-2015 at 05:51

A blank sample was prepared containing only 5 mL deionized water.

One drop of 3M Nitric acid was added to the blank sample.

One drop of 3% Hydrogen peroxide was added to the blank sample.

4 drops of Potassium hexacyanoferrate (II) were added to the blank sample.

After 12 minutes it was noted that the same yellow color had developed as in the Fe3+ containing samples.

14e_Blank+(4)Prussiate+12min.jpg - 375kB

A further control sample was prepared by dissolving a few crystals of FeSO4 in deionized water and a 5 mL aliquot was withdrawn and placed in a vial.

2 drops of 0.2M Potassium hexacyanoferrate (II) were added to the FeSO4 solution and a blue precipitate formed immediately.

One drop each of 3M Nitric acid and 3% Hydrogen peroxide were added to the FeSO4 solution and the blue color of the precipitate intensified. At the same time the quantity of precipitate increased.

15C_FeSO4+Prussiate+Nitric+H2O2.jpg - 342kB

Given the above results, is it possible that the hexacyanoferrate itself is being oxidized to yield the yellow coloration?

Furthermore, how did the blue precipitate form in the FeSO4 sample when only Fe2+ should have been present? (This question is to be taken lightly since I am not confident in the purity of the FeSO4 used, it was a sample of practical grade industrial material.)

[Edited on 29-6-2015 by agent_entropy]

woelen - 29-6-2015 at 05:53

Quote: Originally posted by agent_entropy

I had to be out of the lab for awhile while my first son was born.

<-- shameless bragging
[...]

VERY good reason of being out of the lab! Congratulations!

The blue precipitate in all your samples does not surprise me. In your experiment you make hexacyanoferrate(III), which in solution is unstable and forms blue insoluble complexes. For this reason you need to perform the test within a few hours, otherwise all samples turn blue and you cannot distinguish anymore between the samples.

Interesting to see that your results are very similar with the smaller amount of H2O2 and the extra ferrocyanide. So, my theory of having all ferrocyanide being oxidized to ferricyanide was wrong. It just was a try to make the test more sensitive, but you already have maximum sensitivity.

agent_entropy - 29-6-2015 at 06:05

Thanks! I'm very excited.

I'm also curious about the disappearing color in the thiocyanate samples. I realize the thiocyanate complex is less stable than the ferricyanide, but I would have expected there to be an equilibrium where it re-forms.

Per wikipedia: FeSCN2+ + OH− → FeOH2+ + SCN−

I guess that makes sense. But they were acidic solutions, there shouldn't be much OH- in there.

EDIT:
Nevermind. Same Wiki page, "Thiocyanatoiron decomposes with exposure to light."

[Edited on 29-6-2015 by agent_entropy]

blogfast25 - 29-6-2015 at 07:13

Quote: Originally posted by agent_entropy

EDIT:
Nevermind. Same Wiki page, "Thiocyanatoiron decomposes with exposure to light."

Congrats with the baby. I'll 'wet his head' tonight!

Aha! Light decomposing 'thiocyanatoiron'? That would explain it.

That would also be easy to prevent: a strong excess of SCN-(aq) should mean most Fe3+ is always complexed or at least during measurement/detection, unless the complex Fe(OH)2+ is much stronger than FeSCN2+ but then low pH (low [OH-]) needs to be taken into account too.

Incidentally, FeSO4 will nearly always test positive for Fe(+3) due to partial oxidation and the sensitivity of the 'Prussian Blue' test. 'Good quality' Mohr's Salt - (NH4)2Fe(SO4)2 - is to be preferred to FeSO4 as an Fe(+2) standard because its acidity slows down oxidation of Fe(+2) to Fe(+3). Mohr's salt is easy to prepare and easy to recrystallise for better purity.

[Edited on 29-6-2015 by blogfast25]

aga - 29-6-2015 at 14:47

Quote: Originally posted by agent_entropy

I had to be out of the lab for awhile while my first son was born.

<-- shameless bragging

Congratulations on all fronts !

Take time out from the lab to Be There for your Wife and Son : you will not regret it.

Chemistry can wait.

AJKOER - 30-6-2015 at 05:06

After reading the following found on Atomistry.com on FeCl3 (link: http://iron.atomistry.com/ferric_chloride.html ):

"Alcoholic solutions of ferric chloride are reduced by light, which acts, not as a catalyst, but as a generator of the necessary chemical energy. Ferrous chloride, hydrogen chloride, and formaldehyde are the primary products of the reaction."

It occurred to me that based on photo-fenton chemistry (as some background, see for example, "Fundamental Mechanistic Studies of the Photo-Fenton Reaction for the Degradation of Organic Pollutants" at https://www.google.com/url?sa=t&source=web&rct=j&... ), perhaps an auxilary test based on a non-visual sense as it would be hard to compete with the excellent prior suggested procedures anyway. That is, I am suggesting a test based on the sense of smell relating to formaldehyde formation.

My take on the reaction path: the formation of the methyl radical and subsequent reattack by another hydroxyl radical:

CH3OH + HO. → CH3. + H2O
CH3. + HO. → CH2O + H2O

I know, this test obviously 'stinks', but some may find it useful nevertheless. Note, formaldehyde has been described as being toxic, allergenic, and carcinogenic, so take appropriate safety precautions to limit exposure.
------------------------------------------------

Note, the comments on Atomistry.com are actually extracts from the reputedly authoritative sources over the years. However, some of the material may be quite dated as is the case, I suspect, with the extract above. I would not be surprised if the reaction, however, was discovered owing to the evident formaldehyde formation.

[Edited on 30-6-2015 by AJKOER]

AJKOER - 1-7-2015 at 18:40

Quote: Originally posted by AJKOER

.....

My take on the reaction path: the formation of the methyl radical and subsequent reattack by another hydroxyl radical:

CH3OH + HO. → CH3. + H2O
CH3. + HO. → CH2O + H2O

.....

Actually, I should have posted:

My take on the reaction path: the formation of the methoxy radical and subsequent reattack by another hydroxyl radical:

CH3OH + HO. → CH3O. + H2O
CH3O. + HO. → CH2O + H2O

Here is a confirming source, "Low Temperature Kinetics of the CH3OH + OH Reaction", published in J Phys Chem A. 2014 Apr 17; 118(15): 2693–2701., in particular pages 2 and 17, link: http://www.ncbi.nlm.nih.gov/pmc/articles/PMC4190665/

[Edited on 2-7-2015 by AJKOER]