| Pages:
1
2 |
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
|
|
reaction chloric acid with hydrogen peroxide
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
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
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
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
|
|
(8)HOCl <==> (2)ClO2 + (3)Cl2 + (4)H2O
(2)ClO2 + H2O <==> HClO2 + HClO3
HClO2 + (2)H2O2 --> (2)H2O + HCl + (2)O2
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
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
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 |
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
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
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
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
|
|
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
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)
|
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
Hazard to Others
Posts: 298
Registered: 20-5-2011
Location: New Jersey
Member Is Offline
Mood: No Mood
|
|
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]
hibernating...
|
|
|
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. 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.
|
|
|
| Pages:
1
2 |
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
