Sciencemadness Discussion Board

Hydrochloric acid hydrogen peroxide mix

denatured - 15-9-2010 at 10:58

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 - 15-9-2010 at 11:13

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 - 15-9-2010 at 14:01

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 - 15-9-2010 at 14:07

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.

Eclectic - 15-9-2010 at 15:48

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 - 15-9-2010 at 18:50

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 - 15-9-2010 at 22:05

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 - 15-9-2010 at 22:17

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 - 15-9-2010 at 22:47

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 - 16-9-2010 at 07:37

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.

hissingnoise - 16-9-2010 at 07:56

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?

Formation of Chlorine

AndersHoveland - 8-7-2011 at 17:40

"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]

AJKOER - 8-12-2011 at 13:42

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 - 10-12-2011 at 18:19

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 - 11-12-2011 at 07:34

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 - 11-12-2011 at 15:16

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 - 11-12-2011 at 16:55

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 - 11-12-2011 at 19:07

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 - 11-12-2011 at 21:10

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.



Quote: Originally posted by Bezaleel  

What is being formed here? Is it Cl2 that I smell, or ClO2?


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 - 12-12-2011 at 05:00

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.

AJKOER - 12-12-2011 at 15:51

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 - 13-12-2011 at 18:49

It then seems that especially
Quote: Originally posted by AJKOER  
At 60° another reaction : ClO2=1/2Cl2+O2, sets in.

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 - 14-12-2011 at 09:53

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 - 14-12-2011 at 16:11

Quote: Originally posted by Bezaleel  
What about the possibly evolving perchlorate? How could I reduce that to chloride?


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 - 27-12-2011 at 16:47

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.

reaction chloric acid with hydrogen peroxide

AndersHoveland - 28-12-2011 at 00:09

Quote: Originally posted by AJKOER  
Came across a related observation on the HCl/H2O2 reaction in Mellor.
"E. Lenssen found that hydrogen chloride gives oxygen and the free halogen or chloric acid and water."


The reason for the formation of chloric acid probably relates to the fact that chloric acid does not react with hydrogen peroxide until the temperature reaches 80degC, at which point it is reduced to HCl and Cl2, with the liberation of O2.
"Oxidation and Reduction with Hydrogen Peroxide", Wilder D. Bancroft, Nelson F. Murphy, J. Phys. Chem., 1935, 39 (3), pp 377–398

here is another reference, repeated in this thread again:
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)

Yes, just to confirm, the first reference gives the reaction temperature at 80 degrees and the second reference at 70 degrees.

[Edited on 28-12-2011 by AndersHoveland]

AJKOER - 28-12-2011 at 14:25

AndersHoveland:

I agree with prior comment that on a net basis only HCl and H202 acting together may seem to inefficiently produce HOCl (or Chloric acid for that matter). However, I believe it is revealing to look at the net reaction in possible stages:

Stage 1: HCl + H2O2 = HOCl + H2O

Stage 2: HCl + HOCl = Cl2 + H2O
-------------------------------------------------
Net: 2 HCl + H2O2 = Cl2 + 2 H2O

Note, Stage I is the reaction as cited by Watt's so there is some foundation here, and Stage 2 necessarily follows upon accepting the cited Net reaction.

So even though no net HOCl is apparently produced above, there could still be significant amount of HOCl created as an intermediary. This is may be important looking at an ionic version of the second stage assuming little or no ionization for HOCl:

Stage 2: H(+) + Cl(-) + HOCl = Cl2 + H2O

So in the presence of a reactive metal, for example, the removal of H(+) means the reaction equilibrium moves to the left and more HOCl may be available. Note, there are several preparation methods for aqueous HOCl using this technique (example, adding CaCO3 or HgO or CuO to Chlorine water and filtering).

This suggested analysis could be significant as even low concentration of HOCl can be highly reactive. If, in fact, a high concentration is present, then disproportionation of the HOCl to HClO3 is even more likely.











[Edited on 28-12-2011 by AJKOER]

AndersHoveland - 28-12-2011 at 14:34

(8)HOCl <==> (2)ClO2 + (3)Cl2 + (4)H2O

(2)ClO2 + H2O <==> HClO2 + HClO3

HClO2 + (2)H2O2 --> (2)H2O + HCl + (2)O2

AJKOER - 28-12-2011 at 15:42

AndersHoveland:

Thanks for the reactions.

What caught my eye is a possible Net reaction taking the chemical equations as you have written:

8 HOCl + 2 H2O2 --> 3 Cl2 + 5 H2O + HClO3 + HCl + 2 O2

However, the created HCl could react with more HOCl to produce more Cl2 and H2O:

9 HOCl + 2 H2O2 --> 4 Cl2 + 6 H2O + HClO3 + 2 O2

So any created (or available) HOCl, per the supplied equations, would result in some Chloric acid together with Chlorine and Oxygen gases, as has been reported.

AndersHoveland - 28-12-2011 at 15:47

Quote: Originally posted by AJKOER  


9 HOCl + 2 H2O2 --> 4 Cl2 + 6 H2O + HClO3 + 2 O2

So any created (or available) HOCl, per the supplied equations, would result in some Chloric acid together with Chlorine and Oxygen gases, as has been reported.


Yes, I had this thought after I made the post, but I did not want to go to the trouble of explaining, because these types of reactions can be complicated. The reaction rate between H2O2 and HOCl is significantly faster than the equilibrium between HCl and HOCl. The H2O2, in fact, will tend to reduce any aqueous chlorine to hydrochloric acid. So there is probably not a single ideal equation to describe the reaction. The HCl will actually tend to react as a reducing agent towards the chloric acid. So likely chloric acid cannot be simultaneously formed with hydrochloric acid without the hydrochloric acid bein oxidized to chlorine, which is typically not favorable.

In summary, most of the HOCl will be reduced by the H2O2 to HCl, with a lesser quantity will be oxidized to HClO3 and Cl2. But the HCl will reduce any of the HClO3 that forms. So the only thing that will result from the reaction is chlorine. Excess H2O2 would further reduce the Cl2 to HCl.

The fact that chloric acid can only be obtained as a product when very dilute H2O2 is added to an excess of aqueous Cl2 can potentially give some clues to the chemistry.

Apparently, the Cl2 first hydrolyses with water into HCl and HOCl. The HCl is then oxidizes to another portion of HOCl. The HOCl can then shift its equilibrium towards the formation of ClO2, and additionally Cl2. As soon as the ClO2 forms, the final fate of the reaction is determined, because the ClO2 can hydrolyse into HClO2 and HClO3, the former of which will be reduced, either directly by the H2O2. The likely reason the H2O2 must be so dilute is so that the HClO will have time to shift its equilibrium before it is immediately reduced.

So two competing reactions likely take place, which could be represented, although perhaps somewhat misleadingly, as:

Cl2 + H2O2 --> (2)HCl + O2

Cl2 + (6)H2O2 --> (2)HClO3 + (5)H2O

The latter reaction must predominate under such specific reaction conditions, because apparently there is not enough HCl to oxidize all of the HClO3.

HClO3 + (5)HCl --> (3)Cl2 + (3)H2O

Or essentially, a small portion of chlorine is oxidized to chloric acid, but the oxidation is very inefficient, and the main reaction is simply the catalytic decomposition of the hydrogen peroxide.

To note again, the reaction under such specific conditions, when very dilute H2O2 is added to an excess of Cl2 water, favors different products than when H2O2 is simply reacted with an equivalent of Cl2 water.


I am just going to state now that I am not completely sure of all the details of this type of reaction. The chemistry and equilibriums are fairly complex. The literature seems to suggest that HCl and chlorate exist in equilibrium with chloride and chlorine. And it is very unclear about the equilibrium between HOCl and ClO2/Cl2. There must be an equilibrium, but many texts appear not to make any mention of it, despite the fact that it should be apparent.

If someone else can find some more references (such as the reaction between KClO3 and HCl, that would be helpful.

[Edited on 29-12-2011 by AndersHoveland]

AJKOER - 28-12-2011 at 18:03

AndersHoveland:

I think an important, but perhaps difficult, operational question is what does a freshly prepared solution expected to behave like?

While a final product of an 'old' HCl/H2O2 solution may be closest to Chlorine water, my speculation on a fresh solution undergoing a reaction with a metal, for example, may be somewhat different. That is, the solution temperature is likely raised (ClO2 starting to decompose, and dissolved/created Cl2 and some O2 is leaving the solution), pH is higher per the metal reaction (so there is more HOCl which may be disproportionating into HClO3).

So my speculation is HCl/Cl2/HOCl/O2/HClO3, in declining order of presence.

[Edited on 29-12-2011 by AJKOER]

AndersHoveland - 28-12-2011 at 18:17

As I have previously stated, I have already tried reacting 30% HCl with 30% H2O2. While there is no doubt oxidation and reduction, back and forth, of the HCl, the net reaction is essentially just the gradual decomposition of H2O2. But the solution does turn a yellowish-greensish tinge, and a small amount of chlorine is given off, but most of the gas is just oxygen. Apparently the HOCl (or Cl2) is reduced by the H2O2 at a faster rate than the HCl is oxidized. Left on its own, the solution continues to gradually bubble of between around 1.5 to 3 hours.

Such a solution can rapidly dissolve copper metal, with vigorous bubbling.

Reaction of aqueous chlorine water with hydrogen peroxide, under more typical conditions than described in the previous post, tend to just reduce the chlorine to hydrochloric acid, with the liberation of oxygen.

'Old' HCl/H2O2, at least how I prepared it, is not really any different than when it is freshly prepared, other than the fact that much of the H2O2 has already been decomposed. But the reaction between HCl and H2O2 is obviously dependant on the reactant ratio and concentration.

I suspect that higher pH is detrimental to any potential formation of HClO3.

The reason H2O2 is typically unable to oxidize chloride ions is because it acts to reduce the intermediate oxidation products back to chloride faster than they form.

[Edited on 29-12-2011 by AndersHoveland]

AJKOER - 9-1-2012 at 00:53

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)



It occurred to me that in the case of someone using an HC/H2O2 etching solution on a metal causing an exothermic reaction, for example, that the following reaction might take place as well:

4 HCl + O2 --Heat--> 2 H2O + 2 Cl2 <----> 2 HCl + 2 HOCl

occurring in both concentrated (?) and dilute solutions with a heat source as the reaction of O2 upon HCl is endothermic.

So the products of the second reaction quoted from above, upon netting out the HCl formed under suitable conditions, could be expressed as:

2 HCl + O2 --Heat--> 2 HOCl

implying that in the presence of any sufficiently vigorous reaction, the possible direct oxidation of at least some of the HCl to HClO. Note, the smaller molar concentration of either HCl or H2O2 would limit the HOCl creation.


As a basis see Watt's "A dictionary of chemistry and the allied branches of other sciences", Volume 1, pages 907-908:

"HYPOCHLOROUS Acid. HCl0.—This acid may be prepared:
1. From the anhydride, as just mentioned.
2. By passing air saturated with hydrochloric acid through a solution of permanganate of potassium, acidulated with sulphuric acid and heated in a water bath. The distillate is a solution of hypochlorous acid formed by the direct oxidation of hydrochloric acid: HCl + 0 = HClO."

LINK:
http://books.google.com/books?pg=PA908&lpg=PA908&sig...





[Edited on 9-1-2012 by AJKOER]

LanthanumK - 9-1-2012 at 09:31

The other day I mixed 3% hydrogen peroxide with some hydrochloric acid and thought I smelled a faint smell of chlorine gas.

In most cases of metal oxidation by HCl/H2O2, however, it appears that the hydrogen peroxide forms a thin oxide coating on the metal (clean pink copper changes to the normal orange oxide-coated color in hydrogen peroxide) then the acid dissolves the oxide to form the metal salt.

[Edited on 10-1-2012 by LanthanumK]

AndersHoveland - 9-1-2012 at 21:58

Quote: Originally posted by AJKOER  

"HYPOCHLOROUS Acid. HCl0.—This acid may be prepared:
1. From the anhydride, as just mentioned.
2. By passing air saturated with hydrochloric acid through a solution of permanganate of potassium, acidulated with sulphuric acid and heated in a water bath. The distillate is a solution of hypochlorous acid formed by the direct oxidation of hydrochloric acid: HCl + 0 = HClO."


That is a good find. Someone should really make a compilation of all the different routes to preparing solutions of hypochlorous acid.