iodine and sulfite in acidic medium
Wilco Oelen
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?
Can it better be written as
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
bleaches.
If someone has any idea, I would be pleased.
Wilco
PS: The word 'photo' must be replaced by 'foto' if one wants to send a
personal mail to me. The address ph...@woelen.nl may be removed soon,
if it is spammed too much.
Mohammed Farooq
[Experimental Details...]
> 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.
>
You will find much details of the reaction on the internet when you a
search for *iodometric determination of sulfite* eg
www.kodak.com/US/plugins/acrobat/
en/motion/support/processing/h243/ecr1303.pdf this website mentions
that photographic sulfite does contain some other sulfur containing
anions.
(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 analysis of the
mystery ( I never had a practical experience of analysing sulfite with
iodine) and post a update.
Wilco Oelen
> 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.
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.
> 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.
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).
>
> 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?
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.
>
> 2. Colloidal sulfur is not formed in that reaction?
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.
>
> 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).
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.
>
> 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.
That kind of nice things are completely out of reach for me :-(
>
> 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?
I spent a ml or so of my expensive analytical 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.
>
> 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 analysis of the
> mystery ( I never had a practical experience of analysing sulfite with
> iodine) and post a update.
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
update.
Wilco
Mohammed Farooq
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?
Mohammed Farooq
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.
Mohammed Farooq
SOME MORE OBSERVATIONS:
> Hydrogen peroxide does react with these species, but the reaction is
> not visible.
I repeated this thing a bit more carefully, added 33% extra pure drop
of H2O2 to the yellow solution , and the color disappeared.
> 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.
Similary the reaction of *solid* sodium bisulfite with 33% H2O2 was
too violent , lots of steam and slight explosive like sound.
> 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.
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?
> >
> > 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.
> That kind of nice things are completely out of reach for me :-(
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.
beavith
>Hello,
>
>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+,
ok. very good.
>
>with the H+ quickly being neutralized by the slightly alkaline excess
>of sulfite.
exactly! the H+ is neutralized and the soln remains slightly
alkaline. as long as the soln is alkaline, and any H+ is sucked up by
the alkalinity, the reaction will go to the right, under normal
conditions.
>
>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.
this would be expected. you are essentially starting off with the
right side of your redox reaction and expecting it to go backwards.
>
>In the strongly acidic environment, I expected the following reaction
>to occur (again strongly simplified):
>
> I2 + SO2 + 2H2O --> 2I- + [SO4]2- + 4H+.
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.
>
>However, the products of this reaction are all colorless, so something
>else is occurring as well. Maybe there is some equilibrium reaction?
>Can it better be written as
>
> 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.
no. it supports the evolution of SO2 gas in an excess acid environment
>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).
this should be telling you that the acid environment is the factor,
not the particular acid. time for a new hypothesis.
>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.
naturally.
>
>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.
as expected.
>
>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.
you get points for thoroughness
>
>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.
>
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!
beavith
wrote:
slow down there! interesting reference, tho.
ever hear the expression "when you hear hoof beats think horses, not
zebras?"
Mohammed Farooq
>slow down there! interesting reference, tho.
>ever hear the expression "when you hear hoof beats think horses, not
>zebras?"
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.
Wilco Oelen
<<< lots of stuff skipped >>>
>
> 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!
>
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 analysis 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!
Wilco
Wilco Oelen
> > Hydrogen peroxide does react with these species, but the reaction is
> > not visible.
> I repeated this thing a bit more carefully, added 33% extra pure drop
> of H2O2 to the yellow solution , and the color disappeared.
This surprises me. You did not first get a brown color, due to
elemental iodine? I'll try this at home again with more H2O2. Maybe
there is excess H2O2, which oxidizes the iodine further to colorless
iodate? I'll send an update if I have some time to experiment again.
> > 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.
>
> Similary the reaction of *solid* sodium bisulfite with 33% H2O2 was
> too violent , lots of steam and slight explosive like sound.
Yes, concentrated H2O2 can give funny things :-)
A nice suggestion - a little off-topic. Carefully add some HTH calcium
hypochlorite (65% - 70% active chlorine) to some 30% H2O2 in the dark.
A violent reaction and a beautiful red glow! However, please be
careful and don't use much H2O2!
>
> > 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.
>
> 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?
A high concentration of SO2 stabilizes thiosulfate considerably. This
is used in photography in fixers. Having a high bisulfite content in
the fixer allows a fairly acidic solution to be combined with high
concentrations of thiosulfate, without the thiosulfate decomposing to
sulphur. This may explain, why it does not almost immediately become
milky.
Of course, this is no answer to your question, why the liquid does not
become colorless with thiosulfate. Apparently there really is
something, which strongly stabilises tri-iodide in a solution with
acid and SO2, or even oxidizing iodide to iodine, see remark further
below.
> 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.
This is a very nice result (BTW, thanks for all your efforts and
taking this 'riddle' so serious, I really appreciate that). Your
result convinces me that this really is I3(-), it hardly can be
anything else. This makes the situation very interesting. The only
conclusion I can make with this now is that SO2 or some acidic
species, derived from sulfite is an _oxidizing_ agent, oxidizing I(-)
to I2, although the reaction is not driven to completion.
I think this conclusion is right, because with only iodide, sulfite
and hydrochloric acid, the yellow color is observed. Hydrochloric acid
for sure is not oxidizing, so the only thing left is SO2, or an acidic
derivative of sulfite. This is interesting, because normally SO2 is a
fairly strong reductor.
In inexact equation form.
2I(-) + x SO2 + y H(+) <---> I2 + <reduced species of SO2>
Secondary reaction then is:
I2 + I(-) <---> I3(-)
What <reduced species of SO2> is, I don't have a clue. A nice further
point of research. I have the book "Chemistry of the elements",
written by Greenwood. This desribes many oxo-anions of sulphur. I'll
look at that. If I have an idea, I'll feedback to you.
With this result I can explain the lemon yellow color. A much higher
concentration of iodine is taken up by excess SO2, according to the
following counter reaction, which removes iodine.
I2 + SO2 + 2H2O ---> 2I(-) + SO4(2-) + 4H(+)
This was observed already by means of dissolving I2 in acidic sulfite
solution and iodine being extracted from the ligroin.
>
> 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.
I'll see if I can find this in a second hand book shop. In the recent
past I have obtained some very nice old german books on chemistry for
just a few euros. Here in the Netherlands we have a lot of german
literature on the subject. I like these books, because they do not
only give explanations about the principles, but also give detailed
descriptions of a lot of real-life compounds.
Wilco
beavith
>beavith <beav...@netscape.net> wrote in message
><<< lots of stuff skipped >>>
>>
>> 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!
>>
>
>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 analysis of this phenomenon may be necessary.....
>
i appreciate that. its just that during the school year, we see many
kids obviously doing their homework on the usenet.
if you really are doing this out of curiosity, my hat is off to you.
my days of tinkering at the chemistry set are pretty much over, even
though the set is pretty extensive.
>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!
absolutely.
the question that i was focusing on was the original redox reaction,
which was correct. then you change a condition (an important one!)
and get an unexpected result. my contention is that it shouldn't be
unexpected.
i'm curious about the yellow color too. if MF has the opportunity to
post the UV VIS spectra, i'd guess that the yellow color will either
be I2 or a sulfur dioxide soln. does the mixture smell? does it go
clear when you boil it?
>
>Wilco
Mohammed Farooq
>> 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!
>
>> Wilco
>the question that i was focusing on was the original redox reaction,
>which was correct. then you change a condition (an important one!)
>and get an unexpected result. my contention is that it shouldn't be
>unexpected.
Thats why I was surprised.
>i'm curious about the yellow color too. if MF has the opportunity to
>post the UV VIS spectra, i'd guess that the yellow color will either
>be I2 or a sulfur dioxide soln. does the mixture smell? does it go
>clear when you boil it?
I don't know how to upload a image, neither I do have a scanner but I
have described it in one the follow-ups. It will not take more than 15
minutes to obtain a qualitative spectrum if you (beavith: if you are
seriously interested) have a spectrum recorder.
As I originally assumed that the yellow color is due to iodine as
tri-iodide ion, still the yellow color is not decolorized by sodium
thiosulfate indicating this is not tri-iodide, secondly its spectrum
shows a single sharp peak very close to one of the tri-iodide peaks.
Secondly the solution is very stable.
Secondly sulfur dioxide solution in water is colorless, and if SO2
were causing this color, only acidification of sulfite would have
resulted in a yellow color (the acidified solution does smell of SO2)
but this color is not developed in the absence of iodine or iodide. As
Wilco has written all the three are necessary: acid-sulfite-iodide.
Mohammed Farooq
> > I repeated this thing a bit more carefully, added 33% extra pure drop
> > of H2O2 to the yellow solution , and the color disappeared.
> This surprises me. You did not first get a brown color, due to
> elemental iodine? I'll try this at home again with more H2O2. Maybe
> there is excess H2O2, which oxidizes the iodine further to colorless
> iodate?
Yes I do, when iodide is in excess.
> A nice suggestion - a little off-topic. Carefully add some HTH calcium
> hypochlorite (65% - 70% active chlorine) to some 30% H2O2 in the dark.
> A violent reaction and a beautiful red glow! However, please be
> careful and don't use much H2O2!
You are lucky to observe a red glow somehow I was always fascinated by
chemiluminiscent reactions of hydrogen peroxide. This experiment was
tried long time ago without success with sodium hypochlorite using 6%
medicinal H2O2. What does HTH stand for? Did you use solid calcium
hypochlorite or used a kind of its slurry.
> This is a very nice result (BTW, thanks for all your efforts and
> taking this 'riddle' so serious, I really appreciate that). Your
> result convinces me that this really is I3(-), it hardly can be
> anything else.
But why reaction is taking place id HCl and why thiosulfate does not
react with it?
>This makes the situation very interesting. The only
> conclusion I can make with this now is that SO2 or some acidic
> species, derived from sulfite is an _oxidizing_ agent, oxidizing I(-)
> to I2, although the reaction is not driven to completion.
The formation of slight yellow color after adding H2SO4 to KI solution
is that sulfuric acid does oxidize some iodide to iodine. Sulfur
dioxide can not oxidize iodide to iodine under any condition, because
it is a very poweful reducing agent as you have already said.
> What <reduced species of SO2> is, I don't have a clue. A nice further
> point of research. I have the book "Chemistry of the elements",
> written by Greenwood. This desribes many oxo-anions of sulphur. I'll
> look at that. If I have an idea, I'll feedback to you.
Sulfur dioxide is usually reduced to sulfur, but the solution is
pefectly clear. I don't understand this is such a common reaction with
wide application still no one has described the cause of that yellow
color.
> I'll see if I can find this in a second hand book shop. In the recent
> past I have obtained some very nice old german books on chemistry for
> just a few euros. Here in the Netherlands we have a lot of german
> literature on the subject. I like these books, because they do not
> only give explanations about the principles, but also give detailed
> descriptions of a lot of real-life compounds.
I also like german books ( I can read mainly translations of them)
because of they are really comprehensive and start from basic history
of the concept. In the meanwhile I try to self-learn a working
knowledge of German. Unfortunately neither English nor German books
are available easily nor there are any good chemistry books worth
mentioning in the local langauge.
If you have time take a look at
https://www.chem.ubc.ca/faculty/wassell/CHEM415MANUAL/Experiment2/Experiment2.htm.
and the references therein.
Wilco Oelen
anions. The book I have, mentions that the acidic sulfite / SO2 system
is very complex and several species are in solution at the same time,
being:
SO2
HSO3(-)
HOSO2(-)
O2S2O3(2-)
HO2S2O3(-)
HSO3H
HOSO2H
and possibly others
The metabisulfite (O2S2O3(2-))is the interesting species here, because
it has two sulphur atoms in it, one having oxidation state +5 and the
other having oxidation state +3 (I always thought that both S-atoms
have oxidation state +4 in a metabisulfite ion and I always though the
structure to be like O2SOSO2, but this is false, the two S-atoms are
connected directly, without bridging oxygen). This gives it very
interesting redox properties. The book states that the following net
half-reaction can occur at pH equal to 0:
4SO2 + 4H(+) + 6e <--> S4O6(2-) + 2H2O
The redox potential for the half-reaction, going from left to right
equals +0.509 V, hence SO2 surely can act as oxidizer. The real
reaction is much more complex, than the one given above, it is going
through the O2S2O3(2-) species. The reaction above is just a net
equation, without taking into account all intermediates and
multi-sulphur species.
The lower the pH, the easier the reaction proceeds to the right (the
oxidizer needs H(+)). The redox potential depends on pH. For
decreasing pH, the redox potential becomes even larger.
In the same book, the redox potential for the half-reaction
I2 + 2e <--> 2I(-)
is mentioned as +0.535 V. This does not depend on pH, no H(+) is
involved.
This is close to the redox potential of the reaction above. If pH
becomes lower than 0, then the reaction indeed can proceed somewhat
(at least thermodynamically, and according to our observations, there
apparently is a kinetic route for the reaction as well, otherwise we
would not see the yellow).
The book does not relate iodine with the complex SO2 system. The link
between the two is made by me.
This _may_ be an explanation of what we observe. If this theory is
correct, then it should be possible to observe S4O6(2-) in the
solution, but this will be very difficult to establish (at least for
me, because I do all this at home, without lab facilities). The
S4O6(2-) ion is colorless. This ion is the same as the one, which is
formed when thiosulfate is oxidized by iodine, the so-called
tetrathionate ion.
So, summarizing:
Tetrathionate can be created quantitatively from thiosulfate by
oxidizing this with iodine, but (if the above theory is correct) it
can also be created by reducing acidic sulfite/SO2 with iodide, but
certainly not at quantitative yields.
Wilco
PS: The book I used is "Inorganic Chemistry", by Barett Barnet. It is
an
old book of 1953, and I was lucky enough to find it second hand.
Wilco Oelen
>
> > A nice suggestion - a little off-topic. Carefully add some HTH calcium
> > hypochlorite (65% - 70% active chlorine) to some 30% H2O2 in the dark.
> > A violent reaction and a beautiful red glow! However, please be
> > careful and don't use much H2O2!
>
> You are lucky to observe a red glow somehow I was always fascinated by
> chemiluminiscent reactions of hydrogen peroxide. This experiment was
> tried long time ago without success with sodium hypochlorite using 6%
> medicinal H2O2. What does HTH stand for? Did you use solid calcium
> hypochlorite or used a kind of its slurry.
The calcium hypochlorite, used is solid, just add a spatula of this to
a small amount of 30% H2O2. The letters HTH do not have a special
meaning. This is just a brand of chlorine bleach powder, fairly
commonly available in the Netherlands (I think it originally comes
from the USA). The calcium hypochlorite must be real hypochlorite, not
hypochlorite/chloride. The stuff I use contains Ca(OCl)2, with just a
few percents of other stuff (mostly Ca(OH)2, CaCO3 and also
a little CaCl2). It is technical grade stuff, used for cleaning
toilets, floors etc. and it is also used in swimming pools.
>
> > This is a very nice result (BTW, thanks for all your efforts and
> > taking this 'riddle' so serious, I really appreciate that). Your
> > result convinces me that this really is I3(-), it hardly can be
> > anything else.
>
> But why reaction is taking place id HCl and why thiosulfate does not
> react with it?
See posting, just posted before this one: tetrathionate is proposed as
a _possible_ related compound. Just a theory, no proven fact!
The reason that thiosulfate does not react still is not clear to me. I
also would expect that to react.
>
>
> >This makes the situation very interesting. The only
> > conclusion I can make with this now is that SO2 or some acidic
> > species, derived from sulfite is an _oxidizing_ agent, oxidizing I(-)
> > to I2, although the reaction is not driven to completion.
>
> The formation of slight yellow color after adding H2SO4 to KI solution
> is that sulfuric acid does oxidize some iodide to iodine.
Yes, I agree, at least if the concentration is sufficiently high.
Another source of the slight color may be slow oxidation by oxygen
from air.
> Sulfur dioxide can not oxidize iodide to iodine under any condition, because
> it is a very poweful reducing agent as you have already said.
No, I'm not sure anymore of this, see previous posting, as mentioned
already.
>
> Sulfur dioxide is usually reduced to sulfur, but the solution is
> pefectly clear. I don't understand this is such a common reaction with
> wide application still no one has described the cause of that yellow
> color.
This also surprises me. I really think it is strange that nowhere a
treatise on this reaction can be found.
>
> If you have time take a look at
> https://www.chem.ubc.ca/faculty/wassell/CHEM415MANUAL/Experiment2/Experiment2.htm.
> and the references therein.
This is interesting. It also is about tetrathionate. What surprises
me, however, that nowhere on this site mention is made of the specific
reaction, we are studying now. I cannot imagine that these people
never have seen what we have observed, the more so, because they are
working both with iodine and with oxo-sulphur compounds.
Taking everything into account, I still believe that the yellow color
is due to I3(-), primarily, because the spectrum is so convincingly
close to the one, derived from a reference solution.
What still puzzles me, however, is the extreme stability of the yellow
color (as long as the medium remains acid). Even thiosulfate does not
destroy it. I'll try lateron with a large excess of thiosulfate. If a
small amount of thiosulfate is used, then I can imagine that all
thiosulfate is oxidized by the SO2 and its derivatives, by means of
the following pathway:
1) SO2 et al. are converted to tetrathionate and
iodine is formed.
2) iodine converts thiosulfate to tetrathionate, itself being
reduced to iodide.
If this is true, then iodide would be a catalyst for reaction between
thiosulfate and sulfite, both of them being converted to
tetrathionate. This is just some thinking further and of course all
this must be establised (or disproved) by means of experiments. I'll
put some time in this, within the next days and give an update.
Wilco
beavith
wrote:
>Beavith wrote
>>> 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!
>>
>>> Wilco
>
>>the question that i was focusing on was the original redox reaction,
>>which was correct. then you change a condition (an important one!)
>>and get an unexpected result. my contention is that it shouldn't be
>>unexpected.
>Thats why I was surprised.
>
>
>>i'm curious about the yellow color too. if MF has the opportunity to
>>post the UV VIS spectra, i'd guess that the yellow color will either
>>be I2 or a sulfur dioxide soln. does the mixture smell? does it go
>>clear when you boil it?
>
>I don't know how to upload a image, neither I do have a scanner but I
>have described it in one the follow-ups. It will not take more than 15
>minutes to obtain a qualitative spectrum if you (beavith: if you are
>seriously interested) have a spectrum recorder.
i don't have one. sorry
>As I originally assumed that the yellow color is due to iodine as
>tri-iodide ion, still the yellow color is not decolorized by sodium
>thiosulfate indicating this is not tri-iodide,
nor I2
> secondly its spectrum
>shows a single sharp peak very close to one of the tri-iodide peaks.
that implies one species or multiple species that are extremely
similar.
>Secondly the solution is very stable.
>Secondly sulfur dioxide solution in water is colorless,
but not odorless.
> and if SO2
>were causing this color, only acidification of sulfite would have
>resulted in a yellow color (the acidified solution does smell of SO2)
true, but so should a neutral soln. a small amount of alkalinity is
necessary to eliminate the odor.
>but this color is not developed in the absence of iodine or iodide. As
>Wilco has written all the three are necessary: acid-sulfite-iodide.
knife edge equilibrium between the 3 components?
Mohammed Farooq
> color (as long as the medium remains acid). Even thiosulfate does not
> destroy it. I'll try lateron with a large excess of thiosulfate. If a
> small amount of thiosulfate is used, then I can imagine that all
> thiosulfate is oxidized by the SO2 and its derivatives, by means of
> the following pathway:
> 1) SO2 et al. are converted to tetrathionate and
> iodine is formed.
> 2) iodine converts thiosulfate to tetrathionate, itself being
> reduced to iodide.
> If this is true, then iodide would be a catalyst for reaction between
> thiosulfate and sulfite, both of them being converted to
> tetrathionate. This is just some thinking further and of course all
> this must be establised (or disproved) by means of experiments. I'll
> put some time in this, within the next days and give an update.
>
Wilco,
For the first time a seemingly simple reaction has perplexed me to
this extent. We will find the answer to this problem soon, God
willing.
AN IMPORTANT DEVELOPMENT:
I have found a reference that studies the kinetics (and hence
mechanism) of this very reaction ie iodide-sulfite system in a very
famous journal named "Inorganic Chemistry" published by American
Chemical Society, you will be happy to read its title, a surprising
thing is that it is a very recent article, it means this reaction is
by no menas simple and is still the subject of current studies.
TITLE:
"Non-Metal Redox Kinetics: Reactions of Iodine and Triiodide with
Sulfite and Hydrogen Sulfite and the Hydrolysis of Iodosulfate", Yiin,
B. S.; Margerum, D. W. Inorg. Chem. 1990, 29, 1559-1564.
Sadly our library does not have recent issues of the journal
"Inorganic Chemistry" but any good library would have them. Can you
devote some more time to this problem and just have a look at this
paper. I would like to hear more about it, I assume you have access to
good chemical libraries.
Can anybody help by having a *glance* over this article. I am sure
this would have discussed the possible species responsible for the
yellow color formed in this not-so-simple reaction
Wilco Oelen
> > What still puzzles me, however, is the extreme stability of the yellow
> > color (as long as the medium remains acid). Even thiosulfate does not
> > destroy it. I'll try lateron with a large excess of thiosulfate. If a
> > small amount of thiosulfate is used, then I can imagine that all
> > thiosulfate is oxidized by the SO2 and its derivatives, by means of
> > the following pathway:
> > 1) SO2 et al. are converted to tetrathionate and
> > iodine is formed.
> > 2) iodine converts thiosulfate to tetrathionate, itself being
> > reduced to iodide.
> > If this is true, then iodide would be a catalyst for reaction between
> > thiosulfate and sulfite, both of them being converted to
> > tetrathionate. This is just some thinking further and of course all
> > this must be establised (or disproved) by means of experiments. I'll
> > put some time in this, within the next days and give an update.
Even with the large excess the color does not disappear. Now, this
makes me less confident about that the yellow color is caused by
I3(-). With a large excess of thiosulfate, the liquid becomes milky
after a while, but not immediately. This is just the well-known
decomposition of thiosulfate by acid.
I also tried the following:
Prepare a solution of tri-iodide, by dissolving some I2 in a
concentrated KI-solution and very carefully add thiosulfate-solution
to this, until the liquid JUST becomes colorless. Then it contains
only iodide and tetrathionate (besides of course Na(+) and K(+) ions).
Adding hydrochloric acid (10% by weight) to this colorless liquid does
not give any visible change, the liquid remains colorless. Even after
many hours, the liquid remains clear and colorless. Apparently
tetrathionate is quite stable in a strongly acidic medium. The yellow
color definitely is not due to some iodine-tetrathionate complex.
As soon, as I add a spatula of Na2SO3 to the acidified
tetrathionate/iodide solution, it becomes lemon yellow and it remains
clear!
Now I'm totally lost!
> >
>
> Wilco,
>
> For the first time a seemingly simple reaction has perplexed me to
> this extent. We will find the answer to this problem soon, God
> willing.
>
> AN IMPORTANT DEVELOPMENT:
> I have found a reference that studies the kinetics (and hence
> mechanism) of this very reaction ie iodide-sulfite system in a very
> famous journal named "Inorganic Chemistry" published by American
> Chemical Society, you will be happy to read its title, a surprising
> thing is that it is a very recent article, it means this reaction is
> by no menas simple and is still the subject of current studies.
> TITLE:
> "Non-Metal Redox Kinetics: Reactions of Iodine and Triiodide with
> Sulfite and Hydrogen Sulfite and the Hydrolysis of Iodosulfate", Yiin,
> B. S.; Margerum, D. W. Inorg. Chem. 1990, 29, 1559-1564.
>
> Sadly our library does not have recent issues of the journal
> "Inorganic Chemistry" but any good library would have them. Can you
> devote some more time to this problem and just have a look at this
> paper. I would like to hear more about it, I assume you have access to
> good chemical libraries.
Unfortunately I do not have that kind of access. I'm just a hobbyist,
with self-builtup knowledge of chemistry. I use this knowledge for
enhancing my possibilities in my photo-hobby (see e.g. posting on
vanadium toner in this group and in rec.photo.darkroom as well). My
profession is software engineering and I do not have any professional
connection with chemistry. Obtaining chemicals through photography
shops is not that difficult for me, o.t.o.h., obtaining good journals
on the field is almost impossible for me, they simply are way too
expensive for a private person. A friend of mine works in a chemical
manufacturing company, he _may_ have access to that kind of
literature. I'll ask him, but it will take some time and I cannot
promise anything.
>
> Can anybody help by having a *glance* over this article. I am sure
> this would have discussed the possible species responsible for the
> yellow color formed in this not-so-simple reaction
If there is someone out there with access to the above mentioned
article and with some spare time, I would be really appreciated if one
could give a clue.
My experience with chemistry is that even with very simple things
there are really remarkable things. During my hobby time I have come
accross several riddles, of which this sulfite/iodine riddle is just
one example. Each time, when I find such a riddle, I'm really
surprised to see that no one else apparently ever encountered it (or
the results are simply hidden for me, because they are written down in
journals, not accessible for me).
Wilco
Farooq
Both of us are stunned by this observation- the non-reactivity of the
lemon yellow solution with thiosulfate. The spectrum shows a single
peak in the visible region close to one of the peaks of the
tri-iodide solution. The spectrum suggests something close to iodine
or tri-iodide yet it is stable towards tri-iodide.
If you don't mind, I am posting another message requesting someone who
has access to "Inorganic Chemistry" and has some free time to have a
quick scan over this article to see if the authors discuss anything
about the species responsible for this yellow coloration in the
iodine-sulfite -acid system.
I checked the website of American Chemical Society, they do not
provide abstracts of back issues to non-suscribers. What they ask you
is to buy that article for $ 25 that is pretty expensive and out of my
reach. This riddle has become pretty expensive.
Though you have self learned chemistry, it seems to be much much
better than many chemistry graduates. I think the formal courses
involve you in too much of theoretcial side... studying hypothetical
molecules... yet unaware of real examples like this.
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beavith
wrote:
>ph...@woelen.nl (Wilco Oelen) wrote in message
a 1992 follow on article by the same group is titled
"Non-Metal Redox Kinetics: Reactions of Iodine and Triiodide with
Thiosulfate via I2S2O32- and IS2O3- Intermediates", Scheper, W. M.;
Margerum, D. W. Inorg. Chem. 1992, 31, 5466- 5473.
iodine and iodide stabilized thiosulfate anion? hmmm! that'd
correspond to your single vis peak (or an unresolved double peak).
i can't get into the article either. amazingly, Inorganic Chemistry
archives don't seem to be posted to the web. in this day and age, i
find that to be surprising.
Steve Turner
(Farooq) wrote:
>If you don't mind, I am posting another message requesting someone who
>has access to "Inorganic Chemistry" and has some free time to have a
>quick scan over this article [...]
Which article?
Steve Turner
Farooq
Steve Turner[/quote:b86d390199]
I somehow guessed that you were on vacations. Kindly read the related
post "unexplained color in acid-iodide-sulfite system" which
summarizes the discussion of these 21 follow ups.
The article is
Non-Metal Redox Kinetics: Reactions of Iodine and Triiodide with
Sulfite and Hydrogen Sulfite and the Hydrolysis of Iodosulfate", Yiin,
B. S.; Margerum, D. W. Inorg. Chem. 1990, 29, 1559-1564.
Beavith:
The archives are available of every ACS journal are available see for
example
http://pubs3.acs.org/acs/journals/toc.page?incoden=inocaj&indecade=1&involume=29&inissue=8
SNUMBER6
>B. S.; Margerum, D. W. Inorg. Chem. 1990, 29, 1559-1564.
Whatever else Dale is a great guy ...and a well respected kineticist both from
a theoretical and practical standpoint ... or at least he was 35 years ago when
I knew him ... He's gotta be getting close to 80 now ...Hell, I'm 55 now ...
and I thought he was old back then ... but I bet if you talked to him ... he'll
tell you every detail about this research ...
Be seeing you
In the Village
Number 6