borontrichloride
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Cu(II) coordinated with Ascorbic Acid (Vitamin C)
Greetings all,
This is the first time I have tried experimenting at home so I am a rookie in that respect.
I have copper wire and attempting to formulate copper(II) ascorbate which I then plan to crystallise via vapour diffusion.
I have placed the copper wire into a jar and added an excess of 1.3% w/v sodium hypochlorite.
After an hour, the solution went pitch black. I initially think this is due to the copper wire being dirty and the black is essentially the dirt from
the wire. I have also noticed small blue flakes coming off of the copper wire (likely fragments of Cu(II) entering solution).
My question is: if the black is due to dirt, should I decant the solution and then pour in my pre-prepared basic solution of sodium ascorbate? Or
would the Cu(II) ascorbate form anyway?
Thanks guys,
borontrichloride
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Eddygp
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Can hypochlorite oxidise Cu to CuO? If so, it could be a black suspension too.
As for the blue flakes, if you could actually see them then they are probably insoluble (otherwise they'd just go into solution). So maybe carbonate?
Any more info?
there may be bugs in gfind
[ˌɛdidʒiˈpiː] IPA pronunciation for my Username |
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borontrichloride
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When the solution was black there was no carbonate added. The only thing in that jar when the solution was black was copper wire and 1.3% w/v sodium
hypochlorite.
I did clean the copper wire as much as possible before adding the sodium hypochlorite so I decided there couldn't be that much dirt in there. I made
my ascorbic acid solution with sodium bicarbonate in a 1:4 stoichiometry in a separate vessel with just enough water to dissolve the two and then
added it in.
Within 20 minutes the solution turned green.
I am therefore led to conclude that I have formed CuCl2.
Why this happened after adding the sodium ascorbate solution I don't know yet.
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borontrichloride
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In fact now that I think of it the new reagents must have become oxidised by the hypochlorite and dissociated the O-Cl bonds, thus freeing them to
react with Cu(II).
I also consider that given it is more of an aqua green it may very well be copper(II) carbonate which seems more favourable due to the 1:1
stoichiometry.
[Edited on 8-7-2017 by borontrichloride]
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woelen
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Several things are in play at the same time, obscuring what really happens.
Copper metal certainly is oxidized by hypochlorite ion. At high pH this will lead to formation of CuO. At low pH, the hypochlorite will become
hypochlorous acid, which will decompose to Cl2, hydrochloric acid, possibly some oxygen, and water. Copper certainly will be oxidized in that case, to
blue copper(II) ions, or green chlorocuprate(II) ions if a lot of chloride ion is present.
Copper(II) reacts with ascorbid acid. In the present of chloride ions at low pH, this will lead to formation of CuCl. If pH is high, then you get a
beautiful bright yellow/orange precipitate of hydrous copper(I) oxide.
If you want to understand things better, try to isolate your steps. First assure you have a good solution with copper(II). You can make that from
dilute hydrochloric acid, to which copper is added and a little bleach or hydrogen peroxide. You could also buy some copper sulfate (either online or
locally, this chemical is easy to obtain and not expensive).
Once you have a clear solution, containing copper(II) ions, then you can experiment with ascorbid acid. It is interesting to experiment with this at
different pH. Try dissolving some ascorbid acid in a solution of sodium hydroxide, making an alkaline sodium ascorbate solution, and then add a
solution with copper(II) to this. You see formation of a dark olivegreen complex, and quick subsequent formation of copper(I) oxide. At low pH, the
reaction is slower, but with heating you may even be capable of seeing formation of metallic copper. In the presence of chloride, however, your copper
will not go further than copper(I), due to formation of quite strongly bound chloro complexes.
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borontrichloride
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Yes everything you are saying is what I have assumed. When I get the black solution at first I assume it is CuO. This makes sense in the presence of
NaOCl.
What annoys me is that the step beyond this is non-existent. I add in some sodium bicarbonate and almost immediately get a green-blue solution,
suggestive of CuCO₃.
I try to make CuCl₂ by adding in NaCl and it is not soluble in the NaOCl but they are both ionic so I found it confusing. I appreciate that chloride
is likely a weaker field ligand than oxide but if the chloride is in enough excess I was anticipating ion exchange to form CuCl₂. I know there's
something I haven't considered so I am researching. According to literature, CuCl₂ can be formed through addition of HCl to a Cu(II) salt.
So now I am trying to work out how to make HCl.
Of course I can simply buy HCl, H₂SO₄ and other reagents but I am trying to test myself and see if I can form compounds just using household
products!
Unless I am mistaken, salt and vinegar should mix to form HCl and sodium acetate. Just need to work out molar equivalents.
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borontrichloride
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OK so a quick spot of research indicates that HCl being the stronger acid than acetic acid shows that the acetic acid won't deprotonate for HCl
formation but it will go the other way round.
I have some toilet duck which is 2% w/w lactic acid so I am going to research if lactic acid is a strong enough acid to form HCl. However lactic acid
is organic so I am doubtful.
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MrHomeScientist
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Both acids are sold at hardware stores: HCl as a concrete etchant (muriatic acid), and H<sub>2</sub>SO<sub>4</sub> as a drain
unblocker (the ones wrapped in a thick plastic bag). That qualifies as household products to me!
Another way to make HCl(aq) is with salt (NaCl) and concentrated sulfuric acid, which with heating produces HCl(g) that can then be led into water.
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woelen
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It is interesting to make chemicals yourself, but there are a few very basic chemicals which in my opinion are not worthwhile to make yourself. Those
are H2SO4, HCl (10%), NaOH, H2O2 (6%), and NH3 (5%). They are cheap and easy to obtain. H2SO4 can be purchased in e.g. 37% concentration as battery
acid and can be concentrated by boiling down. In some countries you can even buy the concentrated stuff (96% or so). Making these chemicals yourself
is a pain in the ass. This kind of chemicals can only be made economically in bulk quantities at an industrial scale. Starting from these chemicals it
is interesting to make them at higher concentrations. E.g. I made 15% NH3 from 5% NH3 and 17% HCl from 10% HCl.
I myself did make quite a few chemicals myself in small quantities, but only interesting ones which cannot easily be purchased by private persons.
Some of these are:
- Br2
- PBr3
- KIO3
- KBrO3
- KClO3
- KIO4
- HNO3 (90+ %)
- a whole bunch of transition metal complexes
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