| Pages:
1
2 |
denatured
Hazard to Others
Posts: 151
Registered: 7-8-2004
Location: -
Member Is Offline
Mood: HCl 50%?
|
|
Hydrochloric acid hydrogen peroxide mix
If hydrochloric acid and hydrogen peroxide gets mixed, will I get any vapors (HCl? chlorine oxide(s)?)? how stable is that? is there a kind of
synergism in mixing these two reagents? or is it better to use each separately?
Thanks
|
|
|
DougTheMapper
Hazard to Others
Posts: 145
Registered: 20-7-2008
Location: Michigan, USA
Member Is Offline
Mood: Energetic
|
|
In my experience, even adding 38% HCl to c. 40% H2O2 (as needed for the synthesis of MEKP), no reaction occurs. However, I usually do this at <10
Celsius.
|
|
|
kmno4
International Hazard
Posts: 1503
Registered: 1-6-2005
Location: Silly, stupid country
Member Is Offline
Mood: No Mood
|
|
This mixture is not stable and gives off Cl2 gas, sooner or later but it does.
H2O2/HCl mixture is useful chlorinating agent for some purposes.
[Edited on 15-9-2010 by kmno4]
|
|
|
madscientist
National Hazard
Posts: 962
Registered: 19-5-2002
Location: American Midwest
Member Is Offline
Mood: pyrophoric
|
|
Of course HCl fumes will be present - you're using HCl!
As far as "is it better to use each separately," it's hard to say since you didn't say what you're going to use it for - but judging by your apparent
lack of familiarity with basic chemistry, and the probable plan of action (making acetone peroxide etc.), I'd say keeping them unmixed is wise.
I weep at the sight of flaming acetic anhydride.
|
|
|
Eclectic
National Hazard
Posts: 899
Registered: 14-11-2004
Member Is Offline
Mood: Obsessive
|
|
You might get some chlorine, chorine oxides,,,,the mixture can bubble up and from personal experience, destroy carpeting. Excellent for disolving
seminoble metals like copper, nickle, tin
|
|
|
Anders2
Banned
Posts: 39
Registered: 4-9-2010
Member Is Offline
Mood: No Mood
|
|
I have tried this, no chlorine is produced. I can tell you that you certainly will not be able to get any chlorine oxides by this, since HClO2
immediately reacts with H2O2 to form HCl and O2.
I have read somewhere something about H2O2 oxidizing HCl to make Cl2, but I am unsure, I think the solution would have to be fully saturated with HCl
to make things extremely acidic. Normally acidity stabilizes H2O2 against decomposition.
Theoretically, H2O2 and very concentrated H2O2 should form very minute traces of HClO4, during the decomposition of H2O2 to O2. I tried this, but did
not end up with any detectable perchloric acid, or any perchlorate salts.
I also tried using acetic anhydride to dehydrate a mixture of 30% H2O2 and 12%HCl. Slowly pouring the Ac2O in, at first there was no reaction. More
Ac2O added, then the solution slowly started bubbling just a tiny bit, and continued bubling ever so slightly for several minutes despite no
additional Ac2O being added, as if it were a self-sustaining reaction.
Adding a little more Ac2O, the solution began to violently bubble and got extremely hot to the touch, after one minute it began to foam over and
something started boiling out.
I think the Ac2O was reacting with the H2O2.
Intermediate AcOOAc forming and decomposing into CO2 and methyl radicals, which probably initiated a self-sustaining radical cascade reaction. I am
unsure if this would have happened without using HCl.
[Edited on 16-9-2010 by Anders2]
|
|
|
kmno4
International Hazard
Posts: 1503
Registered: 1-6-2005
Location: Silly, stupid country
Member Is Offline
Mood: No Mood
|
|
Something to read:
THE PROPERTIES OF PURE HYDROGEN PEROXIDE. IV. ACTION OF THE HALOGENS AND HALOGEN HYDRIDES
O. Maass and P. G. Hiebert
DOI: 10.1021/ja01667a005
But in 'Summary' one can read:
1. Hydrofluoric acid does not decompose hydrogen peroxide but acts as a stabilizer. The other halogen hydrides cause the decomposition of hydrogen
peroxide at all concentrations.
For concentrations ~30% of HCl and ~30% H2O2 it gives Cl2, below some concentrations only O2 is liberated.
|
|
|
denatured
Hazard to Others
Posts: 151
Registered: 7-8-2004
Location: -
Member Is Offline
Mood: HCl 50%?
|
|
Thanks everyone for your replies. I was thinking of combining two steps into one, but after reading this, I will stick to normal for now and use HCl
-> wash to neutral -> H2O2.
kmno4: Thanks for digging this info (got pdf?), so at concentrations 4% HCl and 10% H2O2, it should be safe enough.
FYI, I'm not making AP or MEKP or other energetics, I'm just making Chitosan.
|
|
|
woelen
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
You can use a mix of H2O2 and HCl for many different purposes, but you should not store such mixes. If you mix e.g. 25% HCl with 6% H2O2 then you will
not be gassed with Cl2, but storing such a mix is dangerous, because of slow release of Cl2 and pressure buildup in the container.
If you mix 30% HCl and 30% H2O2 then you surely get bubbles of Cl2, but also of O2. You even get some oscillating reaction. One moment the mix
strongly fizzles, some time later it hardly fizzles but its green color intensifies, then it strongly fizzles again and the color becomes somewaht
lighter, and so forth, for many cycles.
|
|
|
madscientist
National Hazard
Posts: 962
Registered: 19-5-2002
Location: American Midwest
Member Is Offline
Mood: pyrophoric
|
|
| Quote: | | FYI, I'm not making AP or MEKP or other energetics, I'm just making Chitosan. |
Your post was vague enough I got a bad feeling from it, sorry about that. 
| Quote: |
I also tried using acetic anhydride to dehydrate a mixture of 30% H2O2 and 12%HCl. Slowly pouring the Ac2O in, at first there was no reaction. More
Ac2O added, then the solution slowly started bubbling just a tiny bit, and continued bubling ever so slightly for several minutes despite no
additional Ac2O being added, as if it were a self-sustaining reaction.
Adding a little more Ac2O, the solution began to violently bubble and got extremely hot to the touch, after one minute it began to foam over and
something started boiling out.
I think the Ac2O was reacting with the H2O2.
Intermediate AcOOAc forming and decomposing into CO2 and methyl radicals, which probably initiated a self-sustaining radical cascade reaction. I am
unsure if this would have happened without using HCl. |
Are you sure you understand the mechanism? Ac2O should react with H2O2 to form AcOH and AcOOH, not AcOOAc. AcOOH is not something you want to
"dehydrate."
What's your source of the Ac2O anyway? It sounds like a catalyst could be breaking down the H2O2, resulting in the exothermic reaction you witnessed.
Trace amounts of iron or manganese will do this.
I weep at the sight of flaming acetic anhydride.
|
|
|
hissingnoise
International Hazard
Posts: 3940
Registered: 26-12-2002
Member Is Offline
Mood: Pulverulescent!
|
|
| Quote: | | What's your source of the Ac2O anyway? |
So, madscientist has an interest in this anhydride - who'da thunk it. . .
But Anders2, do you buy your Ac2O or do you prepare it yourself?
|
|
|
AndersHoveland
Hazard to Other Members, due to repeated speculation and posting of untested highly dangerous procedures!
Posts: 1986
Registered: 2-3-2011
Member Is Offline
Mood: No Mood
|
|
Formation of Chlorine
"The catalytic decomposition of hydrogen peroxide in either hydrogen chloride solution or with chlorine has been shown to be closely related to the
two chemical reactions
H2O2 + (2)H(+) + (2)Cl(-) ---> Cl2 + (2)H2O
H2O2 + Cl2 ---> O2 + (2)H(+) + (2)Cl(-)
and is believed to be due to the occurence of these two competing reactions at equal rates." Livingston and Bray, J. American Chem. Society,
Volume 47, p2069 (1925)
In this reaction, there is mostly not any net generation of chlorine resulting from the action of dilute solutions of hydrogen peroxide on
hydrogen chloride, although elemental chlorine appears to be an important intermediate. Basically, either HCl or Cl2 can slowly catalyze the
decomposition of H2O2 (hydrofluoric acid does not catalyze any reaction). A small amount of the HCl is oxidized by the H2O2 when the solutions are
concentrated, but the reaction is very inefficient. Most of the H2O2 is just decomposed. Interestingly, a similar reaction exists between ozone and
chlorine gas. While the main reaction is almost entirely the catalyzed decomposition of ozone into oxygen gas, there are small traces of ClO2, and
even Cl2O6 (which is a red liquid), which form.
When 30% concentrated HCl and 30% H2O2 is used then there develops a slight greenish yellow color and a faint but distinctive odor of chlorine, yet
the gas from the bubbles is still mostly O2. I have conducted this experiment and observed only moderate steady bubbling that persisted for several
hours.
It is mentioned in the literature that chlorine gas is evolved from 30% solutions of HCl and H2O2, although in more dilute solutions, only oxygen is
generated.
"Oxidation of Hydrogen chloride with hydrogen peroxide in aqueous solution" V.I. Skudaev, A.B. Solomonov
[Edited on 9-7-2011 by AndersHoveland]
I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by AndersHoveland  | "The catalytic decomposition of hydrogen peroxide in either hydrogen chloride solution or with chlorine has been shown to be closely related to the
two chemical reactions
H2O2 + (2)H(+) + (2)Cl(-) ---> Cl2 + (2)H2O
H2O2 + Cl2 ---> O2 + (2)H(+) + (2)Cl(-)
and is believed to be due to the occurence of these two competing reactions at equal rates." Livingston and Bray, J. American Chem. Society,
Volume 47, p2069 (1925)
In this reaction, there is mostly not any net generation of chlorine resulting from the action of dilute solutions of hydrogen peroxide on
hydrogen chloride, although elemental chlorine appears to be an important intermediate. Basically, either HCl or Cl2 can slowly catalyze the
decomposition of H2O2 (hydrofluoric acid does not catalyze any reaction). A small amount of the HCl is oxidized by the H2O2 when the solutions are
concentrated, but the reaction is very inefficient. Most of the H2O2 is just decomposed. Interestingly, a similar reaction exists between ozone and
chlorine gas. While the main reaction is almost entirely the catalyzed decomposition of ozone into oxygen gas, there are small traces of ClO2, and
even Cl2O6 (which is a red liquid), which form.
When 30% concentrated HCl and 30% H2O2 is used then there develops a slight greenish yellow color and a faint but distinctive odor of chlorine, yet
the gas from the bubbles is still mostly O2. I have conducted this experiment and observed only moderate steady bubbling that persisted for several
hours.
It is mentioned in the literature that chlorine gas is evolved from 30% solutions of HCl and H2O2, although in more dilute solutions, only oxygen is
generated.
"Oxidation of Hydrogen chloride with hydrogen peroxide in aqueous solution" V.I. Skudaev, A.B. Solomonov
|
I believe the above comments understated the actual creation of HClO, and many of the observations flow from the properties of dilute and concentrated
HClO solutions (like its decomposition into HCl and O2). In particular, note the following videos showing the dissolving of iron in HCl/H2O2:
http://www.youtube.com/watch?v=XUb3DRb8R6w&feature=relat...
Here HCl / H2O2 mixture with various metals (the 3rd is Iron):
http://www.youtube.com/watch?v=rN5ucv31-_I&feature=relat...
My observation is one of a greenish-yellow reactant mixture with most likely with some Chlorine evolution followed by the characteristic reddish-brown
Ferric Chloride solution. This is precisely the reaction products of adding Fe to a HClO solution (namely, chlorine and FeCl3). Also, the fact that
mixing even dilute 5% H2O2 and 5% HCl to form a more stable HClO appears to be a mystery and befuddles many using this etching solution (see, for
example, "Topic: HCL/H2O2 5% reaction with iron..green then brown violent bubbling,heat?" at: http://www.chemicalforums.com/index.php?topic=46826.0
where again a HClO like reaction is reported.
Now, my cited example involve HCl/H2O2 in the presence of Fe(II) ion, which is known to catalyze H2O2 into a highly reactive Fenton reagent, so the
creation of HClO may be limited to the presence of a catalyst. However, Watt's Dictionary Chemistry cites as a general reaction the creation of HClO
from HCl + H2O2 (not in excess). Also, HCl/H2O2 is a widely used/successfully etching solution.
So I guess we have a disagreement.
|
|
|
Neil
National Hazard
Posts: 556
Registered: 19-3-2008
Member Is Offline
Mood: No Mood
|
|
Perhaps the iron changes something in favor of forming something more exciting then O2 and Cl2 but H2O2 is used orally to induce vomiting, something
which would not be done if it formed HClO in the stomach.
H2O2 also converts HClO to HCl + H2O
Oxidation and Reduction with Hydrogen Peroxide
The Interaction of Hydrogen Peroxide and Hypochlorous Acid in Acidic Solutions Containing Chloride Ion
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Neil:
Yes, on the reaction upon adding excess H2O2 to HCl+H2O2. Reference: "Chemistry, inorganic and organic: with experiments" By Charles Loudon Bloxam,
page 182:
"Hypochlorous acid is formed when a weak solution of hydrogen peroxide is added to a large excess of chlorine water; Cl2 + H2O2 = 2HCl0. With an
excess of the peroxide, HClO + H2O2 = HCl + H20 + 02."
Neal's cited paper on the REDOX reaction with H2O2 does note that under suitable alkaline and concentrations, H2O2 will oxidize (so dilute HCl/H2O2
solutions, meaning higher pH, may act differently?). Also, the author carefully avoids impurities (catalyst?).
Now, with respect to the reaction of HCl/H2O2 both at 30% (or, as I have argued under certain concentration/conditions, or in the presence of a
catalyst, may behave to some extent like HClO), Iron may not be the only special case.
In my second cited video with 3 metals, there is a reaction between Cu and fresh HCl/H2O2, that is, the postulated HClO like equivalent. Normally,
there is no visible reaction between Cu and HCl (there is also a YouTube video on this non event for those who wish to see for themselves), but the
three metal video (which includes Cu, Mg and Fe), clearly demonstrates a vigorous reaction on Cu with the liberation of a gas (O2 perhaps) and the
formation of an intense green compound (possibly CuCl2.Cu(OH)2 ?). Here is another video with a better view of the reaction:
http://www.youtube.com/watch?v=l4FmzLEsbd0
With respect to the chemistry involved, the cited reaction between Cl2O (the gaseous anhydride of HClO) and various metals is:
Cl20 + Fe --> FeO + Cl2
Cl20 +2 Cu + H2O --> Cu(OH)2.CuCl2
2 Cl20 + 4 Ag --> 4 AgCl + O2 (gradual)
Cl2O + 2 Hg --> HgO.HgCl2 (very slow reaction)
I am not sure of the source of this extract from my notes.
Now, my take on the particular reaction with HClO and copper:
HClO + 2 Cu --> Cu2O + HCl
2 HClO +2 Cu --> Cu(OH)2.CuCl2
4 HClO + 2 Cu2O --> 2 CuCl2.Cu(OH)2 + O2 (g)
2 HClO --Cu Catalyst--> 2 HCl + O2 (g)
so Copper (II) Oxygen Chloride (which forms green crystalline needles) and oxygen gas are created. One source Mellor, page 271 (link below), confirms
CuCl2.Cu(OH)2 with no mention of oxygen.
"R. Chenevix notes the ready solubility of cupric oxide in chlorine water,...".
"A. J. Balard found that copper filings are partially dissolved by hypochlorous acid, the soln. after standing some time contains cupric chloride, and
deposits a green pulverulent cupric oxychloride."
__________________
Here is a video that includes the reaction of Al + HCl/H2O2:
http://www.youtube.com/watch?v=ZywexSds-1c
If you freeze frame the on the final product, I believe I see a white salt (an insoluble form of Al(OH)3 ) on top. This is characteristic of the
reaction of HClO + Al, which I have performed myself, which creates this insoluble white Al(OH)3 along with Cl2 visible in solution and some O2 gas.
Referring to "A comprehensive treatise on inorganic and theoretical chemistry", Volume 2 By Joseph William Mellor, page 275:
"P. Grouvelle u reported that aluminium hydroxide suspended in water through which chlorine was passed does not go into soln. Z. G. Orioli obtained a
bleaching liquid by the decomposition of a soln. of bleaching powder with aluminium sulphate, and G. Lunge and L. Landolt found that any aluminium
hypochlorite which may be formed immediately decomposes, liberating hypochlorous acid. A. D. White also found that aluminium is slowly attacked by
hypochlorous acid, and the resulting aluminium hypochlorite immediately decomposes into aluminium hydroxide, oxygen, and chlorine."
LINK:
http://books.google.com/books?pg=PA275&lpg=PA275&dq=...
Now in all the reference videos H2O2 is cited as the catalyst. However, it may be that it isn't, but the metals Fe, Cu, Al and others, are in effect
the real catalyst with respect to the formation of HClO.
[Edited on 11-12-2011 by AJKOER]
[Edited on 11-12-2011 by AJKOER]
[Edited on 11-12-2011 by AJKOER]
|
|
|
AndersHoveland
Hazard to Other Members, due to repeated speculation and posting of untested highly dangerous procedures!
Posts: 1986
Registered: 2-3-2011
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by AJKOER |
"Hypochlorous acid is formed when a weak solution of hydrogen peroxide is added to a large excess of chlorine water; Cl2 + H2O2 = 2HCl0. With an
excess of the peroxide, HClO + H2O2 = HCl + H20 + 02."
|
Excellent and informative find. This reaction is very revealing.
I suspect that a very large excess of chlorine water would need to be used for any HOCl to be obtained, because it may be probable that H2O2
reacts much more rapidly towards HOCl than it does towards HCl.
H2O + Cl2 <==> HCl + HOCl
HCl + H2O2 --> HOCl + H2O
HOCl + H2O2 --> HCl + H2O + O2
Indeed, hydrogen peroxide can actually reduce aqueous chlorine to hydrochloric acid. In water the net equation is,
H2O2 + Cl2 --> 2HCl + O2
| Quote: Originally posted by AJKOER |
under suitable alkaline and concentrations, H2O2 will oxidize (so dilute HCl/H2O2 solutions, meaning higher pH, may act differently?).
|
H2O2 apparently becomes a more reactive and stronger oxidizer under alkaline conditions that slowly catalyze its decomposition. I think the reason may
be the intermediate formation of unstable H2O3.
H2O2 <--> HOO[-] + H[+]
HOO[-] + H2O2 --> HOOOH + OH[-]
The concentration of hydrogen peroxide solutions does not affect its oxidizing power. There are, however, several different types of chemicals that
can act as catalysts to enable H2O2 to act as a stronger oxidizer. These chemicals include acetic acid, sodium hydroxide (especially if boiling),
highly concentrated sulfuric acid, and iron ions (which I suspect may be through the intermediate formation of ferrateVI ). For example, a boiling
solution of ammonium hydroxide and H2O2 can oxidize carbon or plastic. H2O2 can very slowly oxidize NH3, especially when made alkaline with Na2CO3 and
boiling.
Typically acidic conditions act to help stabilize H2O2, since it very slowly decomposes in storage, whereas alkaline conditions can cause it to
decompose within minutes. But some common types of acid have chemical interactions with H2O2, or can actually catalyse its decomposition. Both HF and
H3PO4 can act as stabilizers. Interaction of acetic acid with H2O2 can make it slightly more reactive, through the equilibrium formation of
peroxyacetic acid.
| Quote: Originally posted by AJKOER |
Now, with respect to the reaction of HCl/H2O2 both at 30% (or, as I have argued under certain concentration/conditions, or in the presence of a
catalyst, may behave to some extent like HClO), Iron may not be the only special case.
|
The chemistry of H2O2 and its equilibrium HCl or Cl2 is complicated and fairly complex. Oxidations involving HCl/H2O2 could alternatively be
understood through the intermediate formation of chlorine during the catalysed decomposition. Indeed, aqueous chlorine behaves like HOCl in its
oxidizing ability. The ability of iron to act as a catalyst is very different from the catalytic action of HCl.
I have tried 30% H2O2 with 30% HCl and have observed that it can rapidly dissolve copper. Plenty of oxygen appears to be simultaneosly released from
the reaction. Possibly the copper ions are catalyzing the faster decomposition of H2O2 to release O2.
Cu + H2O2 + (2)HCl --> CuCl2 + (2)H2O
| Quote: Originally posted by AJKOER |
With respect to the chemistry involved, the reaction with various metals is:
2 Cl20 + 4 Ag --> 4 AgCl + O2 (gradual)
I am not sure of the source of this extract from my notes.
|
The reaction of hypochlorous acid (HOCl) solutions with silver oxide only liberates oxygen and produces AgCl.
When chlorine reacts with silver oxide diffused in water, a mixture of silver chloride and silver chlorate is formed.
Silver nitrate can react with sodium hypochlorite to form silver chloride and silver I,III oxide, Ag2O2, is formed, both of which are precipitated. An
unknown substance, with bleaching properties, is left behind in the solution. This substance is unstable, and quickly decomposes after several
minutes, leaving behind silver chlorate in the solution, which does not bleach. If sodium hydroxide is added to the bleaching substance, oxygen gas is
evolved.
However, silver hypochlorite may also be formed from the reaction between silver nitrate and sodium hypochlorite, according to the same book.
If a solution of chlorine is added to excess Ag2O, silver hypochlorite can be formed in solution. AgOCl partially decomposes in darkness, or rapidly
if heated above 60degC, into AgCl and AgClO3.
A comprehensive treatise on inorganic and theoretical chemistry, Volume 2 By Joseph William Mellor. p271
The bleaching substance is probably a mix of ClO2 and HOCl. Although it is really rather speculative on my part, I think the reaction might look
something like:
(6)AgNO3 + (6)NaOCl --> (6)NaNO3 + (4)AgCl + Ag2O2 + (2)ClO2
The formation of oxygen from the addition of NaOH to the bleaching substance probably only occurs in the presence of excess AgNO3 still disolved in
solution. If this is not the case, I cannot see any plausible way that any bleaching compound could be produced in the reaction which would react with
NaOH to produce oxygen. One would expect that an excess ratio of AgNO3 had been used, since if there was any excess NaOCl not reacted, then the
investigators would not have been able to determine that there was a new bleaching substance that had been formed (since NaOCl acts as a bleaching
agent itself). It is, for example, known that Ag2O reacts with HOCl to form AgCl and oxygen gas.
Ag2O + (2)HOCl --> (2)AgCl + H2O + O2
| Quote: Originally posted by AJKOER |
If you freeze frame the on the final product, I believe I see a white salt (an insoluble form of Al(OH)3 ) on top. This is characteristic of the
reaction of HClO + Al, which I have performed myself, which creates this insoluble white Al(OH)3 along with Cl2 visible in solution and some O2 gas.
Referring to "A comprehensive treatise on inorganic and theoretical chemistry", Volume 2 By Joseph William Mellor, page 275:
"P. Grouvelle u reported that aluminium hydroxide suspended in water through which chlorine was passed does not go into soln. Z. G. Orioli obtained a
bleaching liquid by the decomposition of a soln. of bleaching powder with aluminium sulphate, and G. Lunge and L. Landolt found that any aluminium
hypochlorite which may be formed immediately decomposes, liberating hypochlorous acid. A. D. White also found that aluminium is slowly attacked by
hypochlorous acid, and the resulting aluminium hypochlorite immediately decomposes into aluminium hydroxide, oxygen, and chlorine."
|
I would suspect such a reaction might be:
Al2(SO4)3 + (6)NaOCl + (6)H2O --> (3)Na2SO4 + Al(OH)3 + (6)HOCl
where the HOCl would likely decompose unless it was extremely dilute.
(8)HOCl --> (4)H2O + (3)Cl2 + (2)ClO2
The fact that Al(OH)3 is mostly unreacted by aqueous chlorine shows how weak of a base it is.
The exact chemistry is too complicated to fully explain in this post, but there are plenty of other threads about the chemistry of H2O2 and chlorine
oxide solution equilibriums in this forum.
For the catalytic effect of iron ions on hydrogen peroxide, where I suggested that the oxidizing power might be through the transient formation of
ferrate(VI) you might see this thread: "Creating free radicals"
http://www.sciencemadness.org/talk/viewthread.php?tid=17866
[Edited on 11-12-2011 by AndersHoveland]
|
|
|
Bezaleel
Hazard to Others
Posts: 444
Registered: 28-2-2009
Member Is Offline
Mood: transitional
|
|
Interesting. I have a concentrated solution of BaCl2, forming crystals. The solution contains a bit of HCl and a bit of H2O2. pH is around 0. In this
case the [Cl-] is very high. I estimate the [H2O2] to be less than 1%.
Something went wrong in the crystallisation, so I redissolved the BaCl2, added a bit of H2O and heated the solution. At about 50C, the solution turned
pale yellow, and I noticed something that smelled of chlorine.
What is being formed here? Is it Cl2 that I smell, or ClO2?
---------------
If I heat a solution of 36% HCl and 37% H2O2 to about 60C, a runaway reaction takes place, where huge amounts of yellow to orange coloured gas evolve,
the solution quickly heating to its boiling temperature.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by Bezaleel | Interesting. I have a concentrated solution of BaCl2, forming crystals. The solution contains a bit of HCl and a bit of H2O2. pH is around 0. In this
case the [Cl-] is very high. I estimate the [H2O2] to be less than 1%.
Something went wrong in the crystallisation, so I redissolved the BaCl2, added a bit of H2O and heated the solution. At about 50C, the solution turned
pale yellow, and I noticed something that smelled of chlorine. |
I would suspect that you formed some HClO that resulted in an unstable Barium hypochlorite, which upon heating, decomposed releasing Cl2.
| Quote: Originally posted by Bezaleel |
If I heat a solution of 36% HCl and 37% H2O2 to about 60C, a runaway reaction takes place, where huge amounts of yellow to orange coloured gas evolve,
the solution quickly heating to its boiling temperature. |
Interestingly, if you formed HClO at this 36% concentration and heating to 60 C, one would expect concentrated HClO solution to disproportionate.
Possible reactions depending on pH, concentration and reactant ratios could include:
3 HClO --> HClO3 + 2 HCl
6 ClO2 + 3 H2O <==> 5 HClO3 + HCl
8 HOCl <==> 4 H2O + 2 ClO2 + 3 Cl2
8 HCl0 --> 2 HCl03 + 6 HCl + 02
so the smell and color should correspond to Cl2, ClO2 and O2 mix.
It is also possible at high concentration that Chloric acid, HClO3, could further disproportionate into Perchloric acid, HClO4. Per Wikipedia:
"Perchloric acid forms an azeotrope with water, consisting of about 72.5% perchloric acid. This form of the acid is stable indefinitely and is
commercially available. Such solutions are hygroscopic. Thus, if left open to the air, concentrated perchloric acid dilutes itself by absorbing water
from the air.
Dehydration of perchloric acid gives the anhydride dichlorine heptoxide, which is even more dangerous"
|
|
|
AndersHoveland
Hazard to Other Members, due to repeated speculation and posting of untested highly dangerous procedures!
Posts: 1986
Registered: 2-3-2011
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by Bezaleel | The solution contains a bit of HCl and a bit of H2O2. pH is around 0. I estimate the [H2O2] to be less than 1%. I redissolved the BaCl2, added a bit
of H2O and heated the solution. At about 50C, the solution turned pale yellow, and I noticed something that smelled of chlorine.
If I heat a solution of 36% HCl and 37% H2O2 to about 60C, a runaway reaction takes place, where huge amounts of yellow to orange coloured gas evolve,
the solution quickly heating to its boiling temperature. |
For some inexplicable reason, the oxidizing power of hydrogen peroxide seems to significantly increase above degrees C. One could infer that cold
hydrogen peroxide cannot directly attack hydrochloric acid or ammonium hydroxide, except through some sort of highly unfavorable equilibrium, thus
only allowing an extremely slow reaction rate. Heating either increases the oxidizing power of hydrogen peroxide or shifts the equilibrium enough to
allow rapid oxidation.
| Quote: |
Oxidizing Power of Alkaline Hydrogen Peroxide
Although alkaline peroxides (such as CaO2) are stable in the absence of water, hydogen peroxide slowly decomposes in aqueous alkaline solution. A
mixture of hydrogen peroxide and ammonium hydroxide (in a 1:3 ratio) acts as a reactive oxidizer, which can attack organic compounds and elemental
carbon. The reaction rate is negligible at room temperature, but when heated to 60°C the reaction becomes vigorous and
self-sustaining. Such solutions are sometimes known as "base piranha". With a 1:1:5 volume ratio of NH4OH, H2O2, and H2O, respectively, the
half-life times of peroxide were 4 hours at 50°C and 40 minutes at 80°C. "Reaction of Ozone and H2O2 in NH4OH Solutions and Their Reaction with
Silicon Wafers" Japanese Journal Applied Physics. 43 (2004) pp. 3335-3339.
Another stabilizer, magnesium hydroxide, inhibits the formation or reactive radicals in alkaline solutions of hydrogen peroxide, interrupting the free
radical chain reactions by catching the superoxide anion radicals. Zeronian SH & Inglesby MK (1995) "Bleaching of cellulose by hydrogen
peroxide". Cellulose 2: 265-272.
|
Chlorine smells more pungent. Chlorine dioxide has a slightly sweet spicy odor to it. This is probably because ClO2 dissolves at a much slower rate in
water than Cl2. Molecules that are more hydrophobic often smell sweeter or more fragrant.
It is almost certainly just chlorine that was coming out in your reaction, although I suppose, theoretically, it may be possible that ClO2 could be
obtained from some H2O2 and a huge excess of dilute chlorine water. I am not entirely sure what the reaction would be between ClO2 and H2O2.
H2O + ClO2 <==> HClO2 + HClO3
HClO2 + (2)H2O2 --> HCl + (2)H2O + (2)O2
Acidified hydrogen peroxide reduces chlorite to chloride, and does not oxidize chlorate. If the acid is concentrated enough, typically
over 40 percent, the chlorate can be reduced by hydrogen peroxide.
source thread: "KCLO4 from H2O2 ?"
http://www.sciencemadness.org/talk/viewthread.php?action=pri...
I would think the reaction might be (broken into two steps for simplicity),
(2)H2O2 + (2)ClO2 --> HClO3 + HCl + H2O + (2)O2
HClO3 + 5HCl --> 3H2O + 3Cl2
Combining the two above gives a net equation of,
(10)H2O2 + (10)ClO2 --> (4)HClO3 + (8)H2O + (3)Cl2 + (10)O2
|
|
|
Ephoton
Hazard to Others
Posts: 463
Registered: 21-7-2005
Member Is Offline
Mood: trying to figure out why I need a dark room retreat when I live in a forest of wattle.
|
|
I think you will find that solubility of chlorine is lower at higher temps.
HCLO2 will also become H20 and Cl when energy is added.
e3500 console login: root
bash-2.05#
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Here is a reference from Mellor, page 288, that mentions when the presence of sunlight, heat, chlorine, chlorides, ClO2 or platinum can accelerate
some of the previously referenced reactions:
"The aq. soln. [referring to ClO2 in water] is fairly stable in darkness; in sunlight, it decomposes rapidly in a few hours; and slowly in diffused
daylight into chloric acid, HCl03, chlorine, oxygen: 6Cl02 + 2H20 = Cl2 + 02 +4HCl03. Some perchloric acid is formed at the cost of the chloric acid:
2HCl03 +02=2HCl04. The presence of chlorides accelerate the rate of decomposition such that a soln. with 0.15 mol. of chlorine dioxide suffered a 2
per cent, decomposition in five weeks in darkness at 0°, while with a normal soln. of chloride, there was a 70 per cent, decomposition. In the
presence of chlorides the reaction is represented: 6Cl02 + 3H20 = 5HCl03 + HCl; the velocity constants follow the relation d[C102]/dt =
—K[Cl02]2[HCl], and accordingly it is inferred that there is a slow reaction : 2Cl02+H20+HCl=2HCl02+H0Cl, followed by a rapid change:
6HCl02+3H0Cl=5HCl03+4HCl. Platinized asbestos also accelerates the reaction like chlorides. In the presence of chlorine, the reaction progresses: Cl02
+ 1/2Cl2 + H20=HCl03 + HCl, with the side reactions: 6Cl02 + 3H20 = 5HCl03 + HCl, and 3Cl2+3H20=HCl03+5HC1. At 60° another reaction : Cl02=1/2Cl2+02,
sets in. Consequently, the decomposition of aq. soln. of chlorine dioxide is very complex, for there are (i) 2Cl02=Cl2+202, which is accelerated by
raising the temp, or exposure to sunlight; (ii) 6Cl02 + 3H20=5HCl03+HCl, which is accelerated by the presence of chlorides or by platinum; (iii)
2Cl02+1/2C12+2H20=2HCl03 +2HCl, which is accelerated by chlorine; (iv) 3Cl2 + 3H20 = HCl03 + 5HCl, which is accelerated by platinum or chlorine
dioxide; and (v) 2Cl2+2H20=4HCl+02, which is accelerated by light."
[Edited on 12-12-2011 by AJKOER]
|
|
|
Bezaleel
Hazard to Others
Posts: 444
Registered: 28-2-2009
Member Is Offline
Mood: transitional
|
|
It then seems that especially
explains the experimentally observed behaviour, namely the solution turning yellow due to the formation of chlorine that remains dissolved.
| Quote: Originally posted by AndersHoveland | Acidified
One gram sodium chlorate was dissolved in 25 cc of hydrogen peroxide (30%) which had previously been acidified wih 1 cc sulfuric acid (specific
gravity 1.82). The solution was boiled for one hour. Soon after the solution had reached he boiling point a yellow gas was evolved. This was at first
thought to be chlorine but more careful examination showed it to be a mixture of chlorine dioxide and chlorine. Analysis showed small traces of
perchlorate had formed. This experiment showed that chlorate, through the action of acidic solutions of hydrogen peroxide, is largely converted to
chloride. A considerable amount of chlorine and chlorine dioxide is evolved at the same time. Acidic solutions of 3% hydrogen peroxide also were shown
to reduce chlorate to chloride.
It has been suggested that in the above reaction the intermediate formation of small amounts of hydrogen chloride interferes with the reaction,
catalytically causing decomposition of the chlorate. Chlorine is known to react with hydrogen peroxide to form hydrochloric acid and oxygen gas. The
hydrochloric acid thus formed would attack the remaining chlorate, the products of the reaction being chlorine and chlorine dioxide, the chlorine then
reacting with more hydrogen peroxide to again form hydrogen chloride.
Reaction of concentrated sulfuric acid with sodium chlorate did not produce any perchlorate, but it has been reported by other sources that
perchlorate is indeed produced. This may be due to different acid concentrations and ratios of reactants.
The reaction between solutions of chloric acid (HClO3) and hydrogen peroxide does not have any appreciable reaction rate until a temperatures above
70degC. (note that perchlorate is not a reaction product in the decomposition reaction, although it may be likely that traces are formed). experiments
conducted by Sand, published in Zelt phys. Chem.,50, 465 (year 1904) |
So, in order to receive a hypochlorate, chlorate free solution, I should boil the BaCl2 solution making sure the excess [HCl] exceeds the [H2O2], and
reflux until its colour has become clear again?
What about the possibly evolving perchlorate? How could I reduce that to chloride?
|
|
|
busukxuan
Harmless
Posts: 5
Registered: 8-12-2011
Member Is Offline
Mood: No Mood
|
|
Two cases, depending on the amount of H2O2 present.
First case:
H2O2 will oxidize HCl and release chlorine, producing chlorine and water. Aqueous chlorine hydrolyzes into HCl and HOCl(hypochlorous acid) in an
equilibrium reaction.
Or... It oxidizes HCl (as well), producing HOCl, which upon dehydration(opposite of the hydrolysis mentioned above), gives chlorine.
I'm not sure which reaction is major.
Second case:
You have too much H2O2 and it oxidizes the remaining HCl to form more HOCl. I'm not sure if the temperature is enough to support this reaction but I
think it is. Oxides of chlorine might also form(I think that would be Cl2O), but they usually hydrolyze immediately to form chlorine oxoacids(HOCl I
think).
[Edited on 14/12/11 by busukxuan]
[Edited on 14/12/11 by busukxuan]
|
|
|
AndersHoveland
Hazard to Other Members, due to repeated speculation and posting of untested highly dangerous procedures!
Posts: 1986
Registered: 2-3-2011
Member Is Offline
Mood: No Mood
|
|
There are some other threads in this forum about that. Perchlorate is very difficult to reduce. Even Zn/HCl will not reduce it. Perchlorate is best
separated out by fractional crystallization, since its ammonium salt has a very low (surprisingly) solubulity.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Came across a related observation on the HCl/H2O2 reaction in Mellor. To quote:
"E. Lenssen found that hydrogen chloride gives oxygen and the free halogen or chloric acid and water."
SOURCE: "A comprehensive treatise on inorganic and theoretical chemistry" by Joseph William Mellor, page 939.
The authors source on the E. Lenssen reference is Jours, prakt. chem., (1), 81, 276, 1860.
|
|
|
| Pages:
1
2 |
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
