Mixture of HCl+H2O2 with transition elements

Hello guys, today i mixed 28% HCl with 30%H2O2. Is this equation ok?

HCl+H2O2=HClO+H2O

Then i put some amount of elemental copper (from wires) and reaction was quite slow, but after some time it become very vigorous (with a lot of heat and bubbling). At the end solution was deep green and after few minutes it become light blue.

2HClO+Cu=Cu(ClO)2+H2

Is Cu(ClO)2 in the solution or CuCl2*3H2O

In the other experiment i put iron nails into the same mixture, reaction was extremely vigurous and the solution was deep orange at first, but after some time it become dark green.
Why is this happened? Is it by oxidation of oxygen from the air?

Oh dear, a lot of questions...

Thanks for answer @woelen, but after standing iron solution became yellow-brownish. Where are all those impurities and their colours?

Copper solution is still blue.

As i said in one my topic about complexes, i am really into that.

So, i mixed solution of sodium-benzoate with 30% H2O2 and nothing happened. Then i mixed it with a mixture of iron and copper which i obtained before (see the previous post)- i thought about linking benzoate rings with two different metals, as metal centers.

At first purple-brownish precipitate occured, but after some time it became to bubble and reaction became vigurous with a lot of bubbling, heat and steam (i think it was a steam). At the end the solution in the glass was intense yellow.

What happened? Did i make some unstable complex? or precipitate was a catslyst of decomposition H2O2?

I also have a photos, but i am still trying to find a way to post it here.
Any help about unploading in U2Uwill be appreciated.

Quote: Originally posted by Xrpdguy

Hello guys, today i mixed 28% HCl with 30%H2O2. Is this equation ok? HCl+H2O2=HClO+H2O Then i put some amount of elemental copper (from wires) and reaction was quite slow, but after some time it become very vigorous (with a lot of heat and bubbling). At the end solution was deep green and after few minutes it become light blue. 2HClO+Cu=Cu(ClO)2+H2 Is Cu(ClO)2 in the solution or CuCl2*3H2O In the other experiment i put iron nails into the same mixture, reaction was extremely vigurous and the solution was deep orange at first, but after some time it become dark green. Why is this happened? Is it by oxidation of oxygen from the air? Oh dear, a lot of questions...

Quote: Originally posted by Rhodanide

Quote: Originally posted by Xrpdguy   Hello guys, today i mixed 28% HCl with 30%H2O2. Is this equation ok? HCl+H2O2=HClO+H2O Then i put some amount of elemental copper (from wires) and reaction was quite slow, but after some time it become very vigorous (with a lot of heat and bubbling). At the end solution was deep green and after few minutes it become light blue. 2HClO+Cu=Cu(ClO)2+H2 Is Cu(ClO)2 in the solution or CuCl2*3H2O In the other experiment i put iron nails into the same mixture, reaction was extremely vigurous and the solution was deep orange at first, but after some time it become dark green. Why is this happened? Is it by oxidation of oxygen from the air? Oh dear, a lot of questions... Hey rpdguy, regarding the color changes that you saw, that's normal. You made CuCl2 by rxn of Cu + H2O2 + HCl. No Hypochlorous acid was made in any appreciable amount, that's not how those two chemicals react. Copper Hypochlorite doesn't exist either, to my knowledge. With the Iron, you made either FeCl2, or Ferrous Chloride, the green color, which was then Oxidized by O2 in the air to FeCl3, which was the orange color. You probably got a LOT of Cl2 coming off when you did this, yes? SO! To recapitulate: Copper + HCl + H2O2 = CuCl2 (Unbalanced & unfinished equation) | Iron + HCl + H2O2 = FeCl3 (unbalanced & unfinished equation) | OR... Fe + 2HCl = FeCl2 + H2. Hope this clears things up! -Rhodanide

At last i made the photos
1. The black one is probably FeO, cause in reaction with diluted HCl it turned green and some gas has been produced.
2. You can see the colours of copper and iron solution.
3.The solution in the blue glass should be magnesium chloride (cause you said that ClO- ions are not present in the H2O2+HCl mixture). I was curious about a speed of the reacutions between earth alcali elements and transition ones, and there was a lot H2O2 so i tried to decompose it fast by adding KMnO4 crystal and it looks like big mess now.
4. In the right vial is the copper solution with few drops of KMnO4 solution (concentrated), just to decompose H2O2. This one doesnt look like the previous. In the left one is 'something'. I tried to do change ions between copper solution and tin-lead alloy. I looked for tin, but after some time, copper was on the bottom so i thought that reaction is over. Then i add conc. solution of NaOH to gain the hydroxide, but precipitation was yellow (as lead-iodide). None of lead and tin doesnt have that precipitate except iodides. After standing it became blue.

Quote: Originally posted by woelen

A yellow precipitate of copper is possible. When copper(II) is reduced to copper(I) in alkaline solution, then you get Cu2O. If this is formed very slowly, then you get nice crystals, which are brick-red. If this is formed more quickly, then you get an orange, or even yellow, precipitate. The solid then is hydrous Cu2O.

But that precipitate is unstable i think, cause after some time it become blue, as you can see in the left vial.
And whats about magnesium solution? I mixed 2ml of 30%H2O2 and 2ml of 18% HCl, then i add a piece of Mg. Reaction was pretty slow.
*Why does reaction of that mixture is faster with transition elements than earth alkaline (e.g. Mg)?*

After reaction i tried to get Mg(OH)2 by adding (drop by drop) NaOH solution (diluted). I spent about 10ml, but there wasn't a precipitate. By adding drops a lot of bubbles were came out, like it decompose H2O2. Also bad idea (and 'short' nerves ) was adding small KMnO4 crystal to the solution.

My question is: Why there wasnt a precipitate? Did i spend a whole NaOH for neutralisation? If yes, what i have neutraised?

I've been doing hobby chemistry for four years now, almost all inorganic,
and I still do not really understand what I'm doing !

I think that the advice from woelen :

" try to keep experiments as simple as possible. So, do not use mixes of (possibly impure) iron salts and copper salts in your first experiments. First try to understand their behaviour as single chemicals. This makes interpretation of results easier."

is excellent.

I have made a few mixtures of multiple reagents, expecting one thing but getting something else - usually a useless mix of waste products..
or, to re-phrase your own comment
"so i tried to XXX by adding YYY and it looks like big mess now."
was a common result for most of my random experiments.

The interaction and preferred oxidation levels of elements is still baffling and amazing to me.
Understanding even what seem to be simple chemical equations is mind-boggling when you consider intermediate reactions/processes.
The good thing about chemistry is that there is always more to learn
(and here is a good place to learn real chemistry)

Quote: Originally posted by Sulaiman

I've been doing hobby chemistry for four years now, almost all inorganic, and I still do not really understand what I'm doing ! I think that the advice from woelen : " try to keep experiments as simple as possible. So, do not use mixes of (possibly impure) iron salts and copper salts in your first experiments. First try to understand their behaviour as single chemicals. This makes interpretation of results easier." is excellent. I have made a few mixtures of multiple reagents, expecting one thing but getting something else - usually a useless mix of waste products.. or, to re-phrase your own comment "so i tried to XXX by adding YYY and it looks like big mess now." was a common result for most of my random experiments. The interaction and preferred oxidation levels of elements is still baffling and amazing to me. Understanding even what seem to be simple chemical equations is mind-boggling when you consider intermediate reactions/processes. The good thing about chemistry is that there is always more to learn (and here is a good place to learn real chemistry)

I can understand you to the end. An inorganic chemistry is my big hobby and love. Haha in the most cases XXX with YYY isnt ZZZ, but XYZXYZYX

I agree with you and @woelen about reactions and their complexion, but after four years on faculty and 8 years in the lab i saw and did a lot of reactions in Organic and Inogranic and its easy to do some project when you have that explained in the books, but stuffs on this forum arent from the books
The most of these things you cannot find in the books or internet. E.g. on Google i cannot find anything about H2O2 and HCl mixture and its behaviour with metals, even my professors arent sure about some parts and intermedieres. This is real science i think.

I like this forum because we can discuss about our interests and experiences. So i am sure we will rock this topic to the end very soon.

Quote: Originally posted by zed

My concern is... Does this H2O2/HCl reaction, liberate Chlorine gas into the environment? Since the reactant in these cases is Cl2 or Cl-.... Is it going to out-gas? A little whiff of HCl is no big deal, a little whiff of Chlorine, can be very unhealthy.

Chlorine is not good to breathe, but it has excellent warning properties. You immediately smell it and it becomes very irritating when it reaches dangerous levels. You then still have plenty of time to get away into fresh air. I would not worry about being poisoned to death by chlorine gas if you work at a small scale (ml's and not tens of ml's of liquid). Just use common sense and try to avoid excessive breathing of the gas. A small whiff of the gas is not a real issue, just don't stick your nose in a beaker full of chlorine gas

Concentrated H2O2, mixed with concentrated HCl certainly gives chlorine gas, but also a lot of oxygen. You get some oscillatory reaction, with burst of vigorous production of oxygen and periods which are much quieter during which there is (slow) production of Cl2.

@Xrpdguy: Several years of lab work is no guarantee of full understanding. I know someone who works in the lab for more than 15 years now as a synthetic chemist (mainly organic), but he follows many 'recipes' and well-known paths to reach the desired goal. Textbooks contain a lot of important and useful knowledge, but that knowledge is limited. Everyone follows the same well-known paths and that's why many other things are not known at all, or somewhere hidden in old and obscure dusty journals, covered by years of time and being forgotten decennia ago. Even very common compounds, like copper, combined with HCl, or copper, combined with thiosulfates, leads to extremely rich and complex chemistry, which is not described in textbooks, but almost certainly is described in some old journal, for instance in the former USSR or pre-WW-II Germany. It is nearly impossible to find such publications, only few people are interested in these things, because they have no direct commercial application.

Because of this complexity I have decided to adhere to a few rules:
- try to use pure chemicals (within reason and budget of course), not the hardware store stuff. On eBay you can find quite a few different chemicals of good purity for decent prices.
- try to isolate effects and do not proceed before one effect is understood (or at least well-observed and described by you)
- when results become vague, due to addition upon addition of chemicals, quit and start over again.
- select chemicals carefully, so that only one ion/molecule is adding to the reaction. Other ions/molecules in the compound should be spectator ions/molecules. E.g. if studying copper chemistry in sulphuric acid, use copper sulfate and not copper chloride, in order to avoid interference of the newly added species. A safe starting point in dilute aqueous chemistry is that sulfate ion and in most cases also nitrate ions are spectator ions.
- be VERY precise in observing. Colors which are just a little bit different from what you expect can mean that you used impure stuff, but it can also mean that you found something really interesting. An example of the latter can be read here: http://woelen.homescience.net/science/chem/riddles/iodide+su... . I noticed (from other experiments) that iodine, when reduced by acidified sulfite never becomes completely colorless, it always remained a little yellow. This puzzled me.

Hello again,
I did an experiment again. Now with P.A. grade of purity.
I used dry and clean Petri dish to make a big area for the reaction (its hard to obeserve it in the beaker or test tube).
Now i mixed 2ml of 16%HCl (p.a.) with 2ml of 30%H2O2. Within a second i felt a little bit of chlorine, but it desapeared very fast. Then i placed three copper wires, very thin and clean. The reaction starts immediately with some tiny bubbles around the wires.

*That was Chlorine or Oxygen in the bubbles?*

The reaction was faster around the middle wire cause it was longer than other ones. In one moment the solution around the wires became dark blue, almost black. ( I think it was just a high concentration of Cu2+ ions around of it), so i stirred the solution and connected all wires. In that moment the reaction became very fast and there was a lot of chlorine.
I also checked for oxygen by stick, but the flame just above the reaction was intense blue and i lost it very soon, so there wasnt an oxygen.

When all of copper was dissolved, the reaction immediately stopped, so i think that copper is react two or three times faster then other reactants. Then i added a big chunk of copper and it started again.

A good way to slow down the reaction is cooling ( i added 5ml of NH4NO3 conc. solution into mixture) or cool it with ice water bath.

*Can we make an equation now?*

Quote: Originally posted by AJKOER

My take on the chemistry along with other pundits, as I cited in my comments above, is apparently a bit off, per very recent research! Per this source: 'Activation of Oxygen and Hydrogen Peroxide by Copper(II) Coupled with Hydroxylamine for Oxidation of Organic Contaminants', by Hongshin Lee†§, Hye-Jin Lee§, Jiwon Seo§, Hyung-Eun Kim§, Yun Kyung Shin‡, Jae-Hong Kim†, and Changha Lee, Department of Chemical and Environmental Engineering, Yale University and Ulsan National Institute of Science and Technology, published very recently in Environ. Sci. Technol., 2016, 50 (15), pp 8231–8238, DOI: 10.1021/acs.est.6b02067 ........ ........ Basically, Cu(l) formed by the action of H2O2 on Cu(ll), reacts with H2O2 to create the highly active Cu(lll) ion: Cu(l)L + H2O2 --> Cu(lll)L' + Products (see drawing at https://pubs.acs.org/doi/abs/10.1021/acs.est.6b02067 )

Thanks for answering but that equations are opening more questions, like " Why is flame above the mixture blue, if its only oxygen released?" Is the flame of burning stick enough powerful to do an excitation of copper ions? (About 300°C that fire has, i think).

Anyway i am also very into growing crystals in gels, especially dendrimers by using electric current. So after this reaction i took a solution and mixed it with gelatine (70 blooms). After standing there wasnt any crystal even by using 7v battery, so i put a 3ml of 5M Sodium-hydroxide soultion and got blue precipitate immediately. After standing on air precipitate became dark blue, almost dark violete.