Sciencemadness Discussion Board - Suitable ferric compound for redox titrations?

Suitable ferric compound for redox titrations?

blogfast25 - 5-5-2008 at 06:24

I want to carry out some redox titrations using Fe3+ as an oxidising agent and I'm wondering what would be the best Fe (III) compound for this purpose, in terms of general stability, deliquescence/loss of crystal water, etc.

The choice would probably be between ferric ammonium alum (dodeca hydrate), ferric chloride (hexa hydrate, PCB etchant) or ferric oxalate (unknown amount of crystal water, available from silverprint.co.uk), because these are the ones I can get my hands on.

Also, a standard to determine strength (precise molarity of a ≈ 0.1 M ferric titrant), preferably something OTC or home made, would be necessary.

Any suggestions are most welcome!

ScienceGeek - 5-5-2008 at 06:55

What are you titrating against?

You don't have to answer if Iron compounds are your only source of an oxidising ion, but usually there are better oxidising agents than Iron(III) and Iron(II).

blogfast25 - 5-5-2008 at 07:03

I'm titrating against Ti3+, there Fe3+ is commonly used.

Ozone - 5-5-2008 at 07:08

It depends upon your needs, specifically pH, concentration and whether or not the counter-ion is tolerated. Also not that the reactivity of the ferric oxidant is effected by what it is ligated to.

FeCl3 is my ferric iron of choice because it is cheap (600lb at 36% m/m was 140 USD, with freight) and has a high aqueous solubility (my oxidations are aqueous). The down side is that the solutions are highly acidic (high HCl equivalence) and this can wreck your reaction. Any pH greater than 4.5 or so leads to the immediate precipitation of colloidal Fe(OH)3. For water treatment, this is good. For your REDOX (particularly of phenolic species) this is bad (you products will be quite intimate with the floc).

Try Ferric sulfate as well. Both the chloride and sulfate are widely used in water treatment (cheap and readily available).

A good assay involves doing your oxidation and sampling an aliquot against o-phenanthroline (Fe2+(o-phenanthroline)3, ABS 510nm). When standardized against Fe2+ (or Fe3+ with added hydroxylamine), this should tell you how much of your Fe3+ was converted to Fe2+ in the process (beware of pseudo catalytic kinetics with phenolic compounds if air is present) and, hence, the extent of reaction (with knowledge of the stoichiometry).

Cheers,

O3

[edit] Oh yes, remember to assay your starting ferric compound as well. In my experience, FeCl3 at ~78,000 mcg/g made under N2 with degassed water contains approximately 0.6-1.4 % Fe2+. This is required for background correction.

[Edited on 5-5-2008 by Ozone]

blogfast25 - 5-5-2008 at 07:23

Thanks, Ozone, but my needs are a lot simple than that: redox titration of Ti3+ in highly acidic conditions with Fe3+ (using SCN- as indicator) is a very robust procedure for most sample situations. My purpose here is to assay thermite made Titanium metal by dissolving in HCl or conc. sulfuric and titrometric determination of the Ti (as Ti3+) with Fe3+/SCN-. This procedure is used the world over at various stages of Titanox production (I used to work in such a plant, but do not recall what the source of Fe3+ used was... it's over 20 years ago)

[Edited on 5-5-2008 by blogfast25]

Ozone - 5-5-2008 at 13:42

In that case, the chloride (which you can get) or sulfate should be ideal. Pick your acid

.

[edit] What Woelen says, below, is true. Solutions do not keep and will eventually precipitate what looks like hydrated ferric hydroxide.

If you use FeCl3 (because it is easy to get), be sure to standardize your material before use. o-phenanthroline both with and without hydroxylamine will give you total iron and iron(II). Make it fresh and standardize it every day.

Good luck,

O3

[Edited on 6-5-2008 by Ozone]

woelen - 5-5-2008 at 22:29

I do not agree about the chloride. FeCl3.6H2O usually is a very unreliable compound, because it is deliquescent and the amount of water in the commercial stuff is highly variable. I have seen the PCB etchant stuff, which was perfectl dry and rather dark brown, and I have seen wet stuff of a mustard-like color, and even solutions of variable concentration.

The salt FeNH4(SO4)2.12H2O would be the best option if you can obtain that.

Ferric oxalate is an even worse choice than ferric chloride, it is not a true oxalate, but a rather messy complex, containing oxalato-aqua coordinated iron(III). Actually, ferric oxalate is not a well-defined compound, similarly like ferric ammonium citrate, which also is not a well defined compound. Thye compound ammonium ferric oxalate is well defined, it better can be described as ammonium trisoxalatoferrate(III) and it is well-crystallized and air-stable. This also could be an option. I'm not sure though how the oxalato ligand affect oxidizing power of the iron(III).

blogfast25 - 6-5-2008 at 04:25

Quote:

The salt FeNH4(SO4)2.12H2O would be the best option if you can obtain that.

The ferric alum was kind of my compound of choice all along and I've found some Google references in which it is described as a source of Fe3+ for redox titrations. So I'll settle for that.

That leaves the problem of a standardising material, preferably an OTC. I might actually go for high purity Titanium to determine the titration factor, as my purpose is to assay various thermite Ti metals and alloys...

Klute - 6-5-2008 at 07:58

To standardize your Fe3+, you could added an iodide in excess, and back titrate the formed I2 with thiosulfate. KI and Na thiosulfate are readibly available. Na thiosulfate is pretty accurate when preparing solutions, but you could standardize it by titrating with an solid iodate, or dichromate.
Good luck!

blogfast25 - 6-5-2008 at 10:54

Thanks, Klute, but that would potentially lead to a 'standardisation cascade'

. But back-titrating KI --> I2, with thiosulphate/starch, I think I'll try that...

JohnWW - 6-5-2008 at 15:31

The trouble with aqueous Fe+++ salt solutions is that they are extensively hydrolysed, except in very strongly acid solutions, or except if comprehensively complexed by a strong sequestering ligand like porphyrin or EDTA. Also, if HCl (or other halo-acid) is used to prevent hydrolysis, the Fe+++ is converted to the FeCl4- anion. The hydrolysis is liable to interfere with attempts to standardize Fe+++ solutions used as oxidants in titration.

blogfast25 - 7-5-2008 at 03:53

I'm not sure whether hydrolysis would play an important part as a contributor to error, especially in acidic conditions and were the ferric solution is used as a titrant (and not the other way round):

Fe(H2O)n3+ + H2O <--->FeOH(H2O)n-12+ + H3O+

Obviously, at pH << 7, this equilibrium shifts strongly to the left.

But when the weakly hydrolised ferric solution hits the excess of Ti3+ in the titrated sample, reacting the Fe (+III) away further pushes this equilibrium to the left. Potential problems, at too high pH, could occur near the end-point.

To be certain, I'll supplement my NH4Fe(SO4

2.12H2O solution with H2SO4 to avoid hydrolysis of the mother solution... Probably 0.1 M Fe3+/0.05 M H2SO4, something like that...

woelen - 7-5-2008 at 04:10

Not 0.05M H2SO4, better is 1M H2SO4. Ti(3+) also has a tendency to hydrolyse and when it flocculates as hydroxide, it is sufficiently strongly reducing to reduce water (giving H2), so be absolutely sure that you remain on the acidic side.

blogfast25 - 7-5-2008 at 06:53

Woelen:

1 M sulfuric is perhaps better as a precaution than 0.05 M but bear in mind that the Ti3+ is held in a very acidic solution, as it's obtained by dissolving Ti (or alloy) in conc. HCl or conc. sulfuric. Still, an excess acid in the titrant solution is probably a good safeguard against premature hydrolysis of the Ti bearing sample solution...

[Edited on 7-5-2008 by blogfast25]

Klute - 7-5-2008 at 08:06

To dissolve Mohr's salt (the Fe2+ alum), which cna be long with pur dH2O, we suspend the salt in 2/3 the amount of water needed (~650mL for 1L of solution for ex), and add conc H2SO4 in portions with stirring until the salt dissolves completly with the heat liberated.. I suppose you could do the same the with Fe3+ ammonium sulfate, except if it is already readibly soluble in water.

12AX7 - 7-5-2008 at 08:52

Quote:

Originally posted by woelen
Not 0.05M H2SO4, better is 1M H2SO4. Ti(3+) also has a tendency to hydrolyse and when it flocculates as hydroxide, it is sufficiently strongly reducing to reduce water (giving H2), so be absolutely sure that you remain on the acidic side.

Not to mention absolute exclusion of air (a purple solution of Ti(3+) quickly (over days) turns colorless when the metal is removed).

blogfast25 - 7-5-2008 at 09:22

The speed at which Ti3+ oxidises to +IV by air oxygen is often overstated, I feel. I have several quite concentrated solutions Ti3+, some have been transferred in open air from one container to another. No signs of oxidation whatsoever, although it's possible that the fairly high Ti3+ concentration masks that...

The trick to titrate it correctly is to add some Al strip prior to titrating, let that dissolve and the nascent hydrogen reduces any TiO2+ back to +III. Then add sodium bicarbonate, the CO2 will flush out remaining hydrogen and provide a protective blanket against air oxygen. This was the standard procedure where I worked in QC/PC all these years ago. Today they may use potentiometric end-point determination but other than that, I doubt if the procedure has been changed much...

blogfast25 - 15-5-2008 at 09:23

Ok, I've chosen ferric alum solution in 1 M HCl (the titrated solution is about 5 M HCl, so sulfuric acid wouldn't be very practical as an acidifier).

But the alum (from SilverPrint.co.uk) didn't dissolve 100 %, there was some cloudiness (looks like Fe(III) (hydr)oxide), so filtration was called for.

For standardising I'll be using Klute's suggestion of back titration with I2/thiosulphate. But it's a long time ago since I've done this. I seem to remember that S2O32- reacts with dilute HCl and forms Sulphur. Should I run the back titration in alkaline conditions or is the oxidising reaction of the thiosulphate with the iodine much faster than the sulfur generating reaction?

A couple of methods I've come across on the net seem to call for pH < 7, not alkalinity...

Klute - 15-5-2008 at 12:01

No, no alkaline conditions or your formed iodine will dismute to iodate and iodide... the acidic conditions aren't a problem, as you correctly guessed the thiosulfate reacts with the i2 much faster than with any acid.
You can use thiodene as an indicator, although it is already pretty clear when the iodine colour disappears.
Once you've added your excess iodide, leave the erlen covered in the dark for 10-15min to leave the iodine the time to completly form and not sublimate out or get degraded.

blogfast25 - 16-5-2008 at 05:46

Thanks, Klute!

That leaves me with one other question. The Na2S2O3 I bought very recently turns out to be anhydrous, not the usual pentahydrate (it's a photo developer's grade). I'm wondering whether it's suitable for this purpose or whether I'm better crystallising it to pentahydrate... The anhydrous powder seems stable, a well-formed powder, not at all clumpy... I haven't dissolved any of it yet though.

Jor - 16-5-2008 at 06:47

Now that we're talking about iron(III)-compounds...
I have 100g of iron(III)sulphate. This has a really low density and is a light brown very dry powder (not crystalline). It dissolves only VERY slowly. So slow, that I'm not patient enough to wait for it to dissolve. Wiki states iron(III)sulphate nonahydrate is very soluble (440g in 100cc of water , @ 20C). Why is mine so insoluble? Do I have an anhydrous salt maybe?

woelen - 16-5-2008 at 08:39

Anhydrous ferric sulfate does dissolve very well in water, but it takes time. It takes several days before all of it has dissolved. A similar thing is true for anhydrous nickel sulfate. I made the anhydrous salt from the hydrated salt. The anhydrous salt does dissolve, but again, it takes days, maybe 1 week.

You have the anhydrous salt.

[Edited on 16-5-08 by woelen]

Klute - 16-5-2008 at 13:50

The anhydrous thiosulfate should be fine, it's just that you can't really be sure of it cocnentration by analytical weighing, as it could be partially hydrated. So unless you standardize the thiosulfate, you better be off either:
-keeping the anhydrous in a dessicator a day prior weighing an ddo it quick
-form the pentahydrate, dry it well enough, and weigh this.

I would go for the totally anhydrous method. Just accord the weight according to the stoechiometry. You can make a rough titration to be sure you are in a tight margine (sp?).

blogfast25 - 17-5-2008 at 03:19

Klute:

Thanks. For now I'll stick with the anh. thiosulphate and perhaps run a few tests drying it at 110 C, 1 h (230 F) to see if there's any weight change at all. I have no proper lab dessicator, I'm afraid, but could always 'concoct something': anhydrous CaCl2 should be available where I live. Drying at RT would definitely be more desirable.

I've also got K permanganate and K dichromate (but no redox indicator for the latter), So there's scope for alternative methods of standardising.

Yesterday I ran a first real titration of an actual Ti3+ sample (thermite Titanium metal dissolved in 32 w% HCl) with the as yet non-standardised 0.1 M ferric alum solution in 1 M HCl using KSCN as an indicator and it went quite well. The end-point determination will take a little practicing as well as a blank titration because the transition from very pale violet to yellowy - orange (dilute FeSCN2+

isn't very sharp.

[Edited on 17-5-2008 by blogfast25]

Klute - 17-5-2008 at 03:31

To use the dichromate, you can add a precisely weighed amoutn of solid dichromate, add excess iodide, and back titrate the formed iodine with thiosulfate. That should be precise enough.
Of course you can do the same with the permanganate, but i think it's less precise to weigh the permanganate.

Keep the dichromate in a custom dessicator afetr keeping it inthe oven a while, to insure it's anhydrous; I often use a tupperware with a bed of CaCl2 covered by a thick layer of toilet paper as a cheap dessicator for general purposes, or a small jar can with a layer of NaOH/KOH or P2O5 mounted by a plastic screen, avoiding any contact with those agressive materials, for more thorough drying conditions.

Glad to hear it worked out nicely!

blogfast25 - 17-5-2008 at 03:50

Quote:

Originally posted by Klute
Keep the dichromate in a custom dessicator afetr keeping it inthe oven a while, to insure it's anhydrous; I often use a tupperware with a bed of CaCl2 covered by a thick layer of toilet paper as a cheap dessicator for general purposes, or a small jar can with a layer of NaOH/KOH or P2O5 mounted by a plastic screen, avoiding any contact with those agressive materials, for more thorough drying conditions.

Glad to hear it worked out nicely!

OK, dichromate it is then! Good idea about the homemade desiccators too: I've got plenty NaOH and tupperware and stuff for perforated plastic support screens: way to go!

As regards it working out, it's going that way but I'm not out of the woods yet: there's of plenty practicing, determination of repeatability, standardising etc before I can call it a success. But progress is being made and that's what matters the most...

I'll certainly report in full here. :cool:

[Edited on 17-5-2008 by blogfast25]

Klute - 18-5-2008 at 03:07

exactly! Titrations are a matter of practice, you get your ways of doing as you do some, especially when it comes to end -point determination, everyone had got his own little thing

There quite alot of information on different titrations, regeants and indicators to use to titrate such and such compound etc etc
This can be very helpfull when trying to determine kinetics of a reaction for example, preparing stock solution of salts, or determining yields of a product in dilute solution etc

blogfast25 - 18-5-2008 at 09:03

Klute,

Yes, practice is essential, especially when the end-point is a bit 'spongy', as is the case here.

One of the problems here is that I use kitchen foil as a source of Al (with HCl the nascent hydrogen reduces any TiO2+ back to Ti3+

which doesn't dissolve completely (due to Si and/or C, I believe), leaving a clouded sample to be titrated and end-point more difficult to establish.

W/o the Al foil, end-point is much sharper, but the titration reads about 2 ml less, so presumably a small amount of Ti is present in the sample solution as TiO2+, rather than Ti3+. A blank titration (no Ti, but Al nonetheless) consumes about 0.5 ml to get the pale orangy colour from diluted FeSCN2+, against the turbid background...

I may have to invest in p.a. Al ribbon or try blank Mg ribbon. Zn granules would do it too but take too long to dissolve.

I've taken your advice on desiccation and now I've got a neat little CaCl2 desiccator from an old pickling jar (you know the ones with a neat red rubber seal?) My anh. thiosulfate and dried dichromate are drying in there.

And I'm waiting receipt of a more accurate jeweller's pocket scale, so I can get up to four significant digits instead of just three.

blogfast25 - 19-5-2008 at 09:25

And now I'm having a little trouble back-titrating the I2 from Fe3+ + 1/2 I- with thiosulphate. Here too, the end-point appears more spongy that I seem to remember. I'm using homemade potato starch solution as an indicator, added near the end-point.

Problem is that after the blue iodine starch complex has cleared, it returns after a few seconds. Adding more thiosulphate, the blue disappears again, only to reappear in seconds. This makes end-point determination quite problematic.

I thought the problem was the starch, so I titrated the next one without it, only to find that the pale yellow iodine solution near end-point regains in colour quickly. This way about a ml of titrant can be added going from almost clear to slightly yellow and so on.

Then I tested the KI. The silverprint.co.uk grade is slightly yellow so may contain some free iodine (will test that later), so I replaced it with a chemset grade that is perfectly white.

To no avail. The iodine keeps reappearing. It's as if the iodide formed during titration is being re-oxidised slightly, back to I2. I'm using distilled water for the thiosulphate solution.

Any ideas, anyone?

Klute - 19-5-2008 at 10:25

Did you leave the iodine at least 10-15min to form? The reaction is autocatalytic IIRC, so it should go fast once all is donbe. 10-15min covered in the black is essential..

blogfast25 - 19-5-2008 at 11:13

I left it for 5 minutes but not longer. There are some reputable resources that recommend titrating the iodine immediately, to avoid evaporation.

But I will try 15 minutes (timed) tomorrow, inside a cupboard.

Klute - 19-5-2008 at 12:53

Just covered or stopper the erlen you use. We always use tappered soints erlenmeyers for theses titrations, stoppered and at least 15min in the dark. If you get the same problem, it could be the starch although I doubt it.

blogfast25 - 20-5-2008 at 04:03

Hi Klute,

Based on just one single titration I would definitely say the waiting period of 15 mins has improved things quite a lot. There is a marked difference in terms of how quickly the colour returns: the end-point solution remains clear for much longer. I think this is workable now. I'll get a few suitably sized bungs for the beer and wine making shop to properly stopper the EM too...

From a previous test (see higher) I know the starch is not at fault.

I also noticed that when I left a half empty burette (stoppered) overnight (I know, I know, shouldn't do that) that the bottom of the thiosulphate solution had gone cloudy due to S formation. That surprises me a little because my distilled water has a pH of about 8. I call it "distilled" but it's really conderser water from a domestic tumble dryer! It's generally speaking much purer than tap water in terms of salt content (almost zero) but the pH does indicate that something has been carried over.

I may have to try commercial deionised water instead but the commercial grades usually contain some anti-oxidant or other...

blogfast25 - 21-5-2008 at 06:07

The back titration of iodine with thiosulfate continues to cause me headaches. End-point determination has improved somewhat by allowing the I2 to form for a specific amount of time (15 min) but end-point is still too spongy for my liking.

I've replaced the potato starch with laundry starch (1 g in 100 ml of boiling AD) and it changes little, if anything. The only thing I can now think of is the AD as a source of problems, as the KI has also been checked. Last two titrations differed by 1.1 ml!

Klute - 21-5-2008 at 10:37

Hum, indeed that might be the problem, becasue thiosulfate solution are stable for months IME, never seen any cloudiness. How concentrated did you make it? 0.1mol/L?
1.1mL difference is indeed very large... Mayeb try deionised water, even if there is a little antioxidants, it might be less problematic than metaliic salts or something...