Hey guys! .I Took a little piece of CuSO4 solid and I added about 5 mL of H2O2 10 vol. and I ended with a deep violet solution (with CuSO4
undissolved). I have no idea what might cause this color, can be any type of complex Cu (II) with peroxo ligands?
Sorry for my English is not very good, I'm from Argentina.Texium - 30-5-2015 at 20:07
I'm curious, what concentration was your peroxide and where did it come from? Because I tried with my 3% and it clearly doesn't cut it.
Also, welcome to the forum. pato_lp - 30-5-2015 at 20:32
I'm curious, what concentration was your peroxide and where did it come from? Because I tried with my 3% and it clearly doesn't cut it.
Also, welcome to the forum.
Hi there!! It was 10 volume Hydrogen peroxide ( for 1 mole of h2o2 you get 10 mole of o2, that concentration unit we use here, maybe its the same than
3%) and i get it from the drugstore( pharmacy). The thing is I do not use a CuSO4 solution, instead I added the peroxide directly into a small crystal
(of cuso4), and deep violet appeared instantly.pato_lp - 30-5-2015 at 20:34
perhaps copper sulphate have some impurity?agent_entropy - 4-6-2015 at 11:25
I tried it with 34% H2O2 and nothing happened.
In your case, I also suspect impurities, if not in the copper sulfate, then in the H2O2 (if it came from a drugstore it likely has some stabilizers
added anyway).
[Edited on 4-6-2015 by agent_entropy]lsn_ - 5-6-2015 at 04:13
May be some impurities in your copper sulhate,i think hydrogen peroxide can't have a reaction with copper sulhate.Loptr - 5-6-2015 at 04:21
I think hydrogen peroxide can have small amounts of phosphoric acid as a stabilizer.Sulaiman - 5-6-2015 at 09:31
I did a quick test, 35% H2O2 + CuSO4 ... exothermic bubbling with a brown deposit on the test tube wall.
The H2O2 is from a reputable supplier, high purity,
the CuSO4 was from a 'Merit' chemistry set, maybe impure?
So I did another experiment using a different batch of CuSO4.
From left to right;
. 35% H2O2 + lots of CuSO4
. 35% H2O2 with a little CuSO4
. 3.5% H2O2 + lots of CuSO4
. 3.5% H2O2 + a little CuSO4
Sorry about the imprecise quantities, just a quick qualitative experiment.
[Edited on 5-6-2015 by Sulaiman]
OOPS!
I only watched the later test for a few minutes,
whilst posting the above the 35% H2O2 reactions must have gone berserk;
now I intend to make a youtube video using known pure CuSO4, with measured quantities.
(the second test was done with CuSO4 from a 'Discovery World' chemistry lab)
(surprising how many people can donate virtually unused 'chemistry sets')
I intend to use my 'proper' supply of CuSO4 after recrystalisation
that's my plan for tomorrow, not sure what my wifes plan is for me tomorrow
[Edited on 5-6-2015 by Sulaiman]blogfast25 - 5-6-2015 at 10:39
Please remember that hydrogen peroxide is 'ambiguous' in terms of it's oxidising/reducing properties. As a general rule (but there are many
exceptions), in alkaline conditions it acts as an oxidiser, see e.g. Cr(+3) to chromate.
But in acid conditions it can be reduced to water/oxygen, see e.g. dichromate to Cr(+3) with H2O2 in acid conditions.
Some species that are 'easily' oxidised like e.g. Fe(+2) to Fe(+3) can be oxidised by H2O2 in either acid or alkaline conditions.
I know of no Cu(+2) peroxo complexes. And as an oxidiser I think H2O2 is far too 'weak' to kick Cu(+2) to Cu(+3), the latter which would immediately
oxidise water anyway.
The khaki green in some of the test tubes is somewhat reminiscent of a Cu(+1)/Cu(+2) mixed valence complex described by woelen and myself (on this
forum). Is some Cu(+2) being reduced by H2O2 to Cu(+1)?
[Edited on 5-6-2015 by blogfast25]Sulaiman - 5-6-2015 at 11:30
I now have a saturated solution CuSO4 filtered and cooling, hopefully giving crystals in the morning.
pato_lp,
thanks for sharing your observations,
such a seemingly simple mixture ......
for a 1yr semi-noob (me) it is very interesting.
blogfast25,
to do a slightly wider experiment, I guess I'll use
35% and 3.5% H2O2 in neutral, acidic and alkaline conditions.
half with stoichiometric excess of CuSO4 and half with defecit.
I suppose H2SO4 for acid, any recommendations for the base ?
(guaranteed no reaction with CuSO4, I do not yet have the required depth of knowledge)
Any recommendations for pH ?
Or has all this been done before? (where?)
[Edited on 5-6-2015 by Sulaiman]blogfast25 - 5-6-2015 at 12:25
Sulaiman:
H2SO4 is just fine. Actual pH value should not be critical.
Remember that above pH 7, Cu(OH)2 will precipitate but that's fine.papaya - 5-6-2015 at 12:38
What I see from photos looks like what happens when you add Cu2+ ammonia complex compound to H2O2, but I also don't remember if pure CuSO4 will react
with peroxide. Are you sure there's no ammonia in there (or any salts of it) ?Sulaiman - 5-6-2015 at 13:10
I am fairly confident of the H2O2 purity, certified food grade from APC
and the glass distilled and polished water was also from APC
I cleaned the glassware well before use.
The main unknown was the purity of the CuSO4 but I doubt there were any ammonia salts in it.
Assuming that my CuSO4 crystallises out in the morning I will be able to rule that out.
I may even try a melting point check of purity, original stock and recrystalised.
(110 C, just at the limit of my thermometer which is very accurate at 0C and 100C)
P.S. I am still looking for a base to increase the pH that will not react with H2O2 or CuSO4 ... suggestions please .....
[Edited on 5-6-2015 by Sulaiman]papaya - 5-6-2015 at 13:21
I may even try a melting point check of purity, original stock and recrystalised.
(110 C, just at the limit of my thermometer which is very accurate at 0C and 100C)
P.S. I am still looking for a base to increase the pH that will not react with H2O2 or CuSO4 ... suggestions please .....
[Edited on 5-6-2015 by Sulaiman]
Melting point, really ? I thought it dehydrates before/during it decomposes (first to anhydrous CuSO4, then to CuO + SO3)! And H2O2 is much LESS
stable under basic conditions by itself.Sulaiman - 5-6-2015 at 13:44
Yes, that was dumb of me, sorry.
I looked it up on Wikipedia but was too hasty;
Melting point 110 °C (230 °F; 383 K) decomposes (·5H2O)[1] <560 °C decomposes[1]
stupid mistake....... must try harder!
Just tested to be sure ... as you all know, it dehydrates rather than dissolves in its own water of crystallisation, which is what I'd assumed
[Edited on 5-6-2015 by Sulaiman]AJKOER - 5-6-2015 at 20:24
Please experiment absence light, which could produce the hydroxyl radical among other things. See pages 282 and especially 283 of "Fundamental
Mechanistic Studies of the Photo-Fenton Reaction for the Degradation of Organic Pollutants" at https://www.google.com/url?sa=t&source=web&rct=j&... where the presence of Phosphoric acid (in the H2O2) could introduce bisulfate into
the reaction mix with CuSO4.
Following the cited paper, I would express some possible reactions of acidified sulfate ions, in the presence of light and H2O2 (a source of hydroxyl
radicals), as follows:
So, some CuS2O8 (Cupric peroxydisulfate), anyone? Note, I would surmise from the above that the hydroxyl radical and H2O2 are largely consumed in the
presence of an acidified sulfate and uv.
------------------------------------------------------
"Many metals are oxidized by persulfate to form soluble metal sulfates, for example, copper:
Cu + S2O8(-2) ——> CuSO4+ SO4(-2)
Under certain circumstances, hydrolysis of the persulfate anion will yield the bisulfate anion and hydrogen peroxide a kinetically faster oxidant
than persulfate:
H(+) + S2O8(-2) + 2H2O ——> 2HSO4(-2)+ H2O2
Another reaction of note is the acid-catalyzed hydrolysis of persulfate to form peroxymonosulfate anion. Fast, high-temperature, acid hydrolysis
followed by thermal quenching will yield solutions of peroxymonosulfate:
H(+) + S2O8(-2) + H2O ——> HSO4(-) + HSO5"
Also the following comments:
"at different pH: Neutral(pH 3 to 7)
S2O8(-2) + H2O ——> 2HSO4(-) + 1/2O2
[Edited on 6-6-2015 by AJKOER]Sulaiman - 7-6-2015 at 13:05
First, using every base that I have I tried in vain to get my copper sulphate solution to pH > 7
then I checked the pH of my (nearly) saturated copper sulphate solution ... 2.2 ...D'oh !
so, nothing new to report in terms of different reactions.
I expect I'll learn more from watching the 'Copper Carnival' competition.
