AndersHoveland
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Sulfur is a stronger oxidizer than Iodine!
I came across this reference that was very surprising to me, and I wanted to share it.
"...the union of hydrogen and sulphur to form gaseous H2S evolves +3.6 units of heat, while 0.8 unit is absorbed in the formation of hydriodic acid
gas. The author has made new experiments upon this point. Dry hydrogen sulfide was inclosed with a little iodine in a sealed tube and heated to 500°,
but no reaction took place. On the other hand, dry hydrogen iodide and sulphur in a similar arrangement reacted immediately, even in the cold, and on
heating to 100° the reaction was complete."
The same article goes on to say that the reverse of this reaction takes place in the presence of water. That when the tube containing the
hydrogen sulfide and iodine was opened under water, it reacted back into sulfur and hydroiodic acid. Yet the reaction could be reversed yet
again, by making the water contain more than 52% hydroiodic acid.
If iodine is dissolved in 52% HI solution, and hydrogen sulfide is passed into the solution, there is no observable reaction.
Reciprocal Displacements between Oxygen, Sulfur, and the Halogens, when combined with Hydrogen.
by Berthelot
Scientific American: Supplement, Volume 7, No. 175, p2791
Apparently hydrogen iodide is a much stronger reducing agent in its anhydrous form. The reactions must be:
(2)HI + S --> I2 + H2S
H2S + I2 + (2)H2O --> S + (2)I- + (2)H3O+
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AndersHoveland
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Sulfur dioxide also oxidizes hydrogen sulfide at room temperature. When SO2 is simultaneosuly bubbled with H2S into water, the water turns a milky
white color, with a slight yellowish tinge.
(2)H2S + SO2 --> (2)H2O + (3)S
or if you prefer
(16)H2S + (8)SO2 --> (16)H2O + (3)S8
Iodine oxidizes aqueous solutions of sulfur dioxide to sulfuric acid. This only works to make dilute sulfuric acid.
H2O + SO2 + I2 --> H2SO4 + (2)HI
Highly concentrated sulfuric acid reacts with hydrogen iodide in essentially the reverse reaction. This is not the entire reaction, because sulfur,
and even some H2S gas, is also produced.
H2SO4 + (2)HI --> H2O + SO2 + I2
Concentrated sulfuric acid will also similarly react with sodium iodide, and will oxidize hydrogen bromide.
[Edited on 29-9-2011 by AndersHoveland]
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woelen
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This is a common effect. I tried myself by adding dry I2 to a bottle with dry SO2. No reaction occurred. As soon as water is added both chemicals
react to zulphuric acid and HI. The reason for this kind of differences in oxidation strength has to do with whether the reaction products are
favourable or not. In a similar way the presence of certain coordinating agents can strongly affect oxidizing power of metal ions (think of oxidizing
power of CuCl2 vs. Oxidizing power of CuSO4).
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AndersHoveland
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Hydrolysis of Sulfuryl Chloride
Woelen, one reaction you may be interested in is reacting dry SO2 with dry Cl2 gas in sunlight (or UV light).
Supposedly SO2Cl2 (sulfuryl chloride) forms.
A modification of this reaction could potentially be used to make anhydrous sulfuric acid, if there was an exact stoichometric ratio of water present.
(just the right quantity of water- not too much or too little)
H2O + SO2 + Cl2 + γ(uv) --> H2SO4 (conc) + (2)HCl (gas)
However, one of the members noted that 98% conc H2SO4 seems to react with HCl to form some chlorine gas (at room temperature). All of the literature I
have found seems to consistantly say that concentrated sulfuric acid can oxidize HI and HBr, but not HCl. But perhaps the HCl can indeed be oxidized
if the concentration of H2SO4 is high enough (98 percent). If this is the case, than the reaction above could only be used to obtain moderately
concentrated sulfuric acid, perhaps 70 or 80 percent concentration.
In other words, what I am saying is that the hydrolysis of pure SO2Cl2 with water might be unable to produce anhydrous H2SO4. The reaction
could be represented by:
(1+x)SO2Cl2 + (2+x)H2O --> (1+x)H2SO4 + (x)H2O + (2)HCl + (x)SO2 + (x)Cl2
where "x" is some fractional proportion. Note that there is water in both the reactants and products, on both sides of the equation; this represents
that the resulting sulfuric will likely contain a proportion of water. This reaction assumes that equal molar ratios of SO2Cl2 and H2O are
reacted with eachother. If excess SO2Cl2 was used, the reaction would be quite different. I suspect that thionyl chloride would result.
(2)SO2Cl2 + H2O --> SOCl2 + H2SO4 + Cl2
To make matters even more complicated, the partial hydrolysis of SO2Cl2 with water might actually look like this:
(3)SO2Cl2 + H2O --> (2)HSO3Cl + SOCl2 + Cl2
where HSO3Cl is chlorosulfuric acid, and SOCl2 is thionyl chloride. The above three reactions assume HCl can act as a reducing agent, so this
speculation may be baseless.
And obviously if excess water was used, the reaction would be straitforward, and only dilute sulfuric acid and hydrochloric acid would result.
SO2Cl2 + H2O --> H2SO4(aq) + (2)HCl(aq)
Since you already have SO2Cl2, this reaction obviously would not be useful to you, but it might be interesting to do. I would be interested to read
about your thoughts on this.
Sulfuryl chloride (SO2Cl2) can also be prepared by reacting sulfur dioxide with chlorine using an activated carbon catalyst at only 25degC.
http://pubs.acs.org/doi/abs/10.1021/ie50172a008
This reaction could potentially be very useful to the members in the energetics section, as a route to make concentrated
sulfuric acid. But to use sulfuryl chloride to make sulfuric acid, it is imparative that we understand the full chemistry going on.
[Edited on 29-9-2011 by AndersHoveland]
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AndersHoveland
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Preparation of SI2 ?
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“Disulfur diiodide, S2I2, can be synthesized by the reaction of S2Cl2 with HI in CCl4 at room temperature, or better, by the reaction of HI with
S2Cl2 in Freon at −78°C. en hyphen deg . In the latter case, it is obtained as a solid, which slowly decomposes to sulfur (especially S6, S7
and S8) and iodine above −30°C. The compound sulfur diiodide does not exist. Although iodine does not react with sulfur, these
elements do form compounds in solution in AsF5/SO2.”
Inorganic chemistry. Egon Wiberg, Nils Wiberg, Arnold Frederick Holleman
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I was thinking about what possible routes there might be to potentially making sulfur diiodide, since this simple compound is
apparently so elusive. I imagine researchers have already attempted all the obvious routes.
What about reacting excess anhydrous hydrogen iodide with SF4 ?
The only fluroide of iodine that can be isolated is iodine pentafluoride. Presumably SF4 would not oxidize I2.
SF4 + (4)HI --> SI2 + I2 + (4)HF
The iodine could potentially ionize in the form of H2F[+] I[-], which could be problematic. Best to use at least 8 equivalents of HI for every 1 SF4.
What do you think would this work?
Other typical routes, such as SH2 with iodine, or SCl2 with HI, are destined for failure, because in the absence of water iodine is surprisingly more
electropositive than sulfur, and vulnerable to oxidation because of its large atomic radius. Indeed, perhaps the hypothetical compound should be
referred to as "diiodine sulfide".
SF6, not entirely inert!
Another interesting reaction is that between SF6 and HI, which only produces elemental iodine.
"Reaction between Sulfur Hexafluoride and Hydrogen Iodide", D. K. Padma, A. R. Vasudeva Murthy
Usually sulfur hexafluoride is very inert, not even reacting with molten metallic sodium under normal conditions!
wikipedia states:
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There is virtually no reaction chemistry for SF6. It does not react with molten sodium...
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It is described as having "inertness resembling nitrogen":
http://pubs.acs.org/doi/abs/10.1021/ie50447a642
I am very surprised. Apparently HI is one of the scarce regents that will react with SF6 at room temperature.
Although apparently SF6 will react with dissolved metallic sodium in anhydrous liquid ammonia at -64degC,
http://pubs.acs.org/doi/abs/10.1021/ic50018a033
[Edited on 7-12-2011 by AndersHoveland]
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madscientist
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SI<sub>2</sub> is probably just not stable.
I weep at the sight of flaming acetic anhydride.
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woelen
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As AndersHoveland states, many people have tried to make binary sulphur-iodine compounds and no one so far succeeded. So, I think that the proposed
method of reacting SF4 with HI does not work. It's a route, which one would not immediately think of, because SF4 most likely will have somewhat
reducing properties (sulphur likes to be in oxidation state +6) and I hardly can believe that it will be capable of oxidizing HI.
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AndersHoveland
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Quote: Originally posted by woelen  | SF4 most likely will have somewhat reducing properties (sulphur likes to be in oxidation state +6) and I hardly can believe that it will be capable of
oxidizing HI.
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SF4 is a strong fluorinating agent. Alcohols are converted into organic fluorides, while carboxyl groups are converted to
trifluoromethyl groups.
http://www.alfa.com/en/docs/FluorinatingAgents.pdf
(scheme 10, page 6)
SF4 can also act as an oxidizing agent,
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Alkylaminosilanes react with sulfur tetrafluoride at ambient temperature with cleavage of the N-Si bond and formation of an N-S bond to give
alkylaminosulfur trifluorides and fluro(trifluoromethyl)silane.
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Methods of Organic Chemistry (Houben-Weyl),
Josef Houben, K. H. Büchel, H. G. Padeken
p394
I think that SF4 would react with SH2 at room temperature, but I cannot find a reference. Certainly sulfur dioxide oxidizes H2S.
I found a reference to iodine pentafluoride reacting with sulfur to form SF4, removing one of the potential doubts about my proposed reaction,
http://pubs.acs.org/doi/abs/10.1021/ja01488a011
[Edited on 8-12-2011 by AndersHoveland]
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