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1 23rd January 00:05
External User
Posts: 1
Default iodine and sulfite in acidic medium


As a photographer hobbyist, with some chemistry knowledge, I came
across the following 'riddle'. Besides being a photo hobbyist, I'm
also interested in chemistry. I'm just curious about what happens
here. I cannot explain it with my knowledge, one of you may be able to
explain it?

I did the following experiments:

Add a granule of iodine to a solution of sodium sulfite. This little
piece of iodine quickly dissolves and the liquid becomes colorless,
just as expected. The reaction, which occurs, I think, is the
following (highly simplified, highlighting main redox effect):

I2 + [SO3]2- + H2O --> 2I- + [SO4]2- + 2H+,

with the H+ quickly being neutralized by the slightly alkaline excess
of sulfite.

I also did this experiment in a fairly acidic environment. I took a
small spatula of sodium sulfite and dissolved this in a few ml of
dilute H2SO4 (appr. 1 mol/l). I added a small piece of iodine and
again this dissolves. To my surprise, the liquid does not become
colorless, it becomes yellow. I first thought that I did not add
sufficient sulfite, but adding more sulfite did not change the color.
Besides that, the liquid has a very pungent odour of SO2, indicating a
large excess amount of sulfite.

In the strongly acidic environment, I expected the following reaction
to occur (again strongly simplified):

I2 + SO2 + 2H2O --> 2I- + [SO4]2- + 4H+.

However, the products of this reaction are all colorless, so something
else is occurring as well. Maybe there is some equilibrium reaction?

I2 + SO2 + 2H2O <--> 2I- + [SO4]2- + 4H+?

Because of lower pH, more H+ is present and this drives the reaction
to the left, such that a clearly visible amount of iodine is present
at equilibrium?

I did two counter-experiments to test the equilibrium hypothesis:
1) Dissolve some potassium iodide (KI) in a solution of sodium
sulfite. This liquid is colorless, as expected. Then add some
dilute sulphuric acid. The liquid becomes yellow at once and
there is a pungent odour of SO2. This supports the hypothesis
of the equilibrium.
2) Dissolve some KI in dilute sulphuric acid. This liquid is
colorless (after several minutes, it turns very pale brown,
but this color is MUCH lighter than the yellow described
above). When some solid sodium sulfite is added, then the
liquid at once turns yellow. Besides this, the smell of SO2
can be observed again. This does not support the hypothesis.
The addition of KI to sulphuric acid should already result
in the yellow color, but the sulfite really is needed to get
the yellow color.

All observations above do not depend on the acid. I also tried all
experiments with dilute hydrochloric acid (appr. 10% commercial acid
from GAMMA, a dutch hardware store, which carries a nice colorless
grade of dilute hydrochloric acid, other stores often have acid with
yellow/green impurities, which are bad for this experiment).
The experiments were also carried out with potassium bisulfite from
another supplier, and the results were consistent with the results,
described above, except, of course that the bisulfite is fairly acidic
on its own already.

In order to test my hypothesis about the equilibrium further, I did
the following two additional experiments:

Make two equal parts of the yellow liquid with excess of sulfite
(strong smell of SO2), each appr. 2 ml large.

Take 1 ml of ligroin (bp. range 40 - 60 C) and add to 2 ml of the
yellow liquid and shake. The ligroin does not become purple, it does
not extract iodine from the aqueous layer.

Take another 1 ml of ligroin and dissolve a small piece of iodine in
this. The ligroin becomes dark purple. Add this to the other half of
the yellow liquid and shake. The ligroin quickly looses its iodine. It
becomes colorless and the aqueous layer remains yellow.

The ligroin experiments do not support the equilibrium hypothesis. I
expected to see at least some purple/pink color in the ligroin, but
not even the faintest pink could be observed in it after shaking a
long time and letting the layers settle again.

So I wonder, what can the yellow stuff be? I'm really surprised by the
results of these experiments. Just simple iodine and sulfite... Does
iodide or iodine form some colored compund with SO2? I never heard or
read about such a thing and a search on Internet did not give any
answer to me. I'm posting these questions, just driven by curiousity
and in order to learn a little more about chemistry. Remember, I'm not
a professional in the field.

The chemicals, used in the experiments are all 'photo grade'
chemicals, obtained at internet shops for raw photography chemicals. I
think, that these are sufficiently pure and that the yellow color is
not due to some impurity. If you have doubts, just check up in your
labs with real reagent grade chemicals!

Please, no discussion about safety on acids, SO2 or the like. I am
aware of the risks of performing chemical experiments and I'm pretty
confident about me knowing what I'm doing. I have worked with nastier
things, like sulfide baths, giving off 'rotten egg' H2S, or dichromate

If someone has any idea, I would be pleased.


PS: The word 'photo' must be replaced by 'foto' if one wants to send a
personal mail to me. The address may be removed soon,
if it is spammed too much.
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2 23rd January 00:05
External User
Posts: 1
Default iodine and sulfite in acidic medium

You will find much details of the reaction on the internet when you a
search for *iodometric determination of sulfite* eg
en/motion/support/processing/h243/ecr1303.pdf this website mentions
that photographic sulfite does contain some other sulfur containing

(I am not a professional too) I tried to repeat your experiment
qualitatively with relatively pure (General Puropse Reagents-GPR )
using sodium bisulfite NaHSO3 , potassium iodide, and extra pure
sulfuric acid since I thought the acid might be the source of
contaminants. As acidified potassium iodide was mixed with sodium
bisulfite solution the mixture immediately acquired a yellow color.
Further addition of iodine solution (brown in color: made by mixing KI
with hydrogen peroxide) to that solution had no effect on the yellow
color and iodine seemed to dissolve in it. In another similiar mixture
addition of hydrogen peroxide had no effect on it indicating that the
yellow color is not due to a reducing species which I was first
thinking of as some sort of (perhaps yellow) oxo-anions of sulfur
mainly pyrosulfite ( synonym: metabisulfite) formed by the reaction of
SO2 and HSO3(-) in equilibrium. So the reaction with hydrogen
peroxide helped to rule out such reducing species in solution.
Secondly I thought due to some impurity , mainly thiosulfate,
colloidal sulfur would have formed when your post was first read but
when your experiment was repeated the solution was perfectly clear.

Do you agree that:
1. The yellow color is not due to some complex oxo-anions of sulfur
(or a hypothetical iodine-SO2 compound) because it has no reducing
property. Also acidificatin of sodium bisulfite solution did not
produce coloration, indicating that yellow color is due to a iodine
containing molecule?

2. Colloidal sulfur is not formed in that reaction?

It is strongly suspected now that the yellow color is due to the
triiodide ion I3(-), since very similiar color is obtained when you
drop very small quantity of solid iodine in potassium iodide solution
and shake for a while ( triodide is known to form by free iodine and
iodide ions).

A UV-Vis spectrum would have immediatly helped to decide if I had time
and quartz cuvettes (for a complete spectrum in the UV region) to
compare the spectrum of a known tri-iodide solution with the yellow
solution so obtained. If they were same then the suspect would indeed
be tri-iodide.

Another thing which is perplexing that even in hydrochloric acid the
same reaction occurs, because HCl is unable to oxidize iodide to free
iodine (which is causing slight yellow coloration in H2SO4)!. I
couldn't repeat your experiment with extra pure HCl . The formation of
sulfur dioxide is explainable here but not the yellow color ie from
where did free iodine come to form yellow triiodide ion?

A suggestion: Try carrying out the experiment quantitatively ie use
every solution of known concentration and use known masses if you have
a nice balance at home, this would help in further ****ysis of the
mystery ( I never had a practical experience of ****ysing sulfite with
iodine) and post a update.
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3 23rd January 00:05
External User
Posts: 1
Default iodine and sulfite in acidic medium

I have some sodium metabisulfite and potassium metabisulfite at home,
and these are white solids. The metabisulfite S2O5(2-) ion is
colorless. In water it goes into equilibrium with bisulfite (H2O +
[S2O5]2- <--> 2[HSO3]-), so this cannot explain the yellow color.
Hydrogen peroxide does react with these species, but the reaction is
not visible. The fact that a reaction occurs, however, can be clearly
deduced, because mixing a solution of metabisulfite with hydrogen
peroxide of sufficient high concentration results in clearly
noticeable warming of the liquid.

Indeed, the solution remains perfectly clear. I kept the yellow
solutions for several hours and they remained perfectly clear and
there was absolutely no precipitate at the bottom.
Another reason for excluding the presence of thiosulfate is that
photo-grade sulfite must be ABSOLUTELY free of thiosulfate. The
sulfite is used in developers and even very small amounts of
thiosulfate would dissolve all silver in the undeveloped
paper/negative, resulting in ruining the undeveloped print. (Ag[+] is
complexed very well by thiosulfate. This is the principle behind
fixer, which removes unexposed Ag[+] from the developed print).

I agree that the yellow color is not due to some oxo-anions of sulfur,
otherwise the yellow color could be observed by careful oxidation of
sulfite with other compounds than iodine. The yellow color is specific
to the iodide/sulfite/acid system. I did some acid/sulfite experiments
with other oxidizing agents (H2O2, KMnO4, Br2), but the resulting
liquid always was colorless, when excess sulfite was used. Especially
the experiment with Br2 is interesting. With Br2 no yellow color is
formed. The Br2 was created by adding a pinch of KBrO3 to an acidified
solution of KBr, such that the liquid becomes orange/brown and a pale
brown vapour of Br2 was above the liquid. At this orange/brown liquid,
an excess of solid sulfite was added, resulting in immediate
disappearance of the orange/brown color and appearance of the smell of
SO2. At this, just a few small crystals of KI were added and
immediately, the colorless liquid turned yellow again!
I'm not sure, whether I can agree with the statement that the yellow
color is not due to some iodine-SO2 or iodide-SO2 complex. I did an
experiment with a drop of H2O2 (10%) added to excess amount of yellow
liquid and I observed formation of a brown cloud, which however,
disappeared on shaking. This can be explained by assuming that iodine
is formed, which however is reduced again, when the liquid is shaken.
The excess of reducing agent (the excess of SO2 in the acidic sulfite
solution) reduces the free iodine again.

I completely agree with this, the liquid remains clear, no milky
appearance at all and after many hours, no white/pale yellow
precipitate at all.

This is an interesting thing. I will think of ways to check whether
this is true or not. I think I3(-) can be detected with starch, also
in acidic environments, I will try this.
Indeed, I3(-) is formed easily. I added a very small piece of iodine
to a solution of potassium iodide. The little piece of iodine did not
dissolve completely, but the liquid was coloured brown. However, I
found that the color is not exactly the same, it is more brown/yellow.
The sulfite/iodide/acid system produces a more bright yellow color. Of
course, I must admit, that this kind of observations is always very
subjective and personal, so I would not add too much value on this observation.

That kind of nice things are completely out of reach for me :-(

I spent a ml or so of my expensive ****ytical grade dilute HCl (2
mol/l) on this 'riddle', but again, yellow stuff, when combining this
with iodide and sulfite. But the yellow stuff only appears when all of
HCl+iodide+sulfite is used. Leave out one of these ingredients, and
you are left with a colorless liquid. I also tried with bromide
instead of iodide, but then the result is a colorless liquid.
Especially the result with HCl really makes me believe that a species
of sulfite or SO2 and iodide is formed, which has the yellow color.
Your hypothesis about I3(-) being formed may also be valid, but then
the only oxidizer which I can imagine is SO2. Normally SO2 acts as a
reductor for iodine, but it is known, that it can act as oxidizer as
well, albeit only in rare cases. May be this is such a rare case? What
the reduced species of SO2 then will be, I have no clue about that.

Unfortunately I do not have an accurate balance (normally,
photographic recipes only need resolutions at the level of a few
grams), but I'll try whether I can derive results from solutions of
known concentrations. If I have any results in this, I'll post an

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4 23rd January 00:05
External User
Posts: 1
Default iodine and sulfite in acidic medium

This is making more interested in solving the mystery behind the
yellow color. This is such a common reaction but unfortunately I am
unable to find specfic mention of yellow coloration during the
reaction. I will see if I get permission to use the a UV-Vis
spectrophotometer to compare the spectrum of tri-iodide with the
yellow color so obtained from this reaction. This will not take more
than 15 minutes if I get permission.

Try diluting the solution of KI and iodine till it matches the color
of the iodine-sulfite-acid mixture. Indeed a concentrated tri-iodide
is dark yellow brown but my dilute tri-iodide solution really matched
the yellow color of this system.

As you say that tri-iodide can be detected by starch, that is true but
we would not be sure which "species" either free iodine I2 (aq) (no
matter how little may be present) or tri-iodide ion which is in
equilibrium with iodine is turning the starch blue.

Can somene else help here?
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5 23rd January 00:05
External User
Posts: 1
Default iodine and sulfite in acidic medium

Just thinking of very well known *clock reaction* called Landolt
reaction,( which uses iodine-sulfite system), the solution changes
color after certain time intervals in a periodic fashion, google it if
you are more interested in this system, in the meantime I will try to
obtain the spectra.
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6 25th January 06:46
External User
Posts: 1
Default iodine and sulfite in acidic medium

I repeated this thing a bit more carefully, added 33% extra pure drop
of H2O2 to the yellow solution , and the color disappeared.

Similary the reaction of *solid* sodium bisulfite with 33% H2O2 was
too violent , lots of steam and slight explosive like sound.

Infact I added sodium thiosulfate to the yellow solution to confirm if
it were tri-iodide ion, the solution should have immeditely
decolorized but nothing happened! What is causing the yellow color now?

Finally took the spectrum of both solutions using distilled water as
*blank*, this subtracts any absorption of light due to water.
Here is the data using Beckmann quartz cuvettes:

The spectrum of tri-iodide has three peaks, a peak near 200 nm is too
intense so it is out of paper, this is not our concern, another far
less than the previous one at 287 nm and another at *347 nm* of
nearly equal intensity.

The spectrum of yellow solution has the same out-of-scale peak at near
200 nm, and surprisingly another peak at *346 nm* !!! Too close to one
of the peak in tri-iodide, indicating that the yellow color is due a
species which is either identical to or very closely related to the
tri-iodide ion.

I do not think now that a complex between sulfur dioxide and elemental
iodine is formed because the absorption maxima of both spectrums are
very very close.

Do you have Holleman and Wiberg's "Inorganic Chemistry" (originally in
German but English translation is available now), My library doesn't,
but I have heard that it is a very comprehensive +2000 pages book,
just check it in a public library, I am sure that book might have
discussed the reaction. I have checked Cotton's "Advanced Inorganic
Chemistry" but could find specific mention of yellow coloration.
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7 25th January 06:46
External User
Posts: 1
Default iodine and sulfite in acidic medium

exactly! the H+ is neutralized and the soln remains slightly
alkaline. as long as the soln is alkaline, and any H+ is ****ed up by
the alkalinity, the reaction will go to the right, under normal conditions.

this would be expected. you are essentially starting off with the
right side of your redox reaction and expecting it to go backwards.

this would be more like
I2 + SO2 + X H+ --> ?

at some point, you'd get preferential release of the SO2 as gas, which you do experimentally.

no. it supports the evolution of SO2 gas in an excess acid environment

this should be telling you that the acid environment is the factor,
not the particular acid. time for a new hypothesis.

you get points for thoroughness

yeah. not a professional. i think i'm doing your homework for you!
tell me this isn't a school lab. for petes sake this is July!
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8 25th January 06:46
External User
Posts: 1
Default iodine and sulfite in acidic medium

On 5 Jul 2004 23:40:58 -0700, (Mohammed Farooq)

slow down there! interesting reference, tho.

ever hear the expression "when you hear hoof beats think horses, not
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9 25th January 06:46
External User
Posts: 1
Default iodine and sulfite in acidic medium

Modify zebra to hinny, Landolt reaction is based on the
iodine-sulfite system and is not totally unrelated!

You have still haven't tried to solve the main question: What is
causing the bright lemon-yellow color?
For basic spectral data, read one of my follow-up...if interested.
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10 25th January 06:47
External User
Posts: 1
Default iodine and sulfite in acidic medium

It looks the signal on this channel is contaminated with some biased
random noise. The source of this random noise may be either arrogance
or ignorance. More ****ysis of this phenomenon may be necessary.....

OK, let's get serious again. I do not say that this is rocket science,
but it sure is not the plain schoolbook chemistry which is involved
here. Also have a look at the very nice investigations of Mohammed
Farooq. Sometimes simple things are really surprising!

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