I am trying to synthesize some copper hydroxide by method of electrolysis of copper. So far i have made two attempt both seemed to produce hydroxide
at first, but as the reaction continued a yellow cloudiness appeared in the electrolyte solution. The first time i though it might be a result from
using a steel cathode but my second try i am using two copper electrodes and the yellowish substance has again appeared in the solution after about an
hour of running the cell. Can anyone tell me what this yellow substance can be and if it will ruin my yield of copper hydroxide thanks. jokull - 12-6-2010 at 11:54
That yellowish substance could be copper chloride complexes. Perhaps the problem is in the quality of water used for your electrolysis.Squall - 12-6-2010 at 11:57
I guess i can try distilled water next and see if that helps right now I am trying a salt bridge but its seems painstakingly slow and i am not sure
what that will yield and i am using distilled water this time
[Edited on 12-6-2010 by Squall]12AX7 - 12-6-2010 at 13:54
What electrolyte?
The yellow to brick-orange precipitate is Cu2O, generally the result when chloride is present.
I am trying to synthesize some copper hydroxide by method of electrolysis of copper.
If I were King - people who post questions like this
would have to state why they choose their method
rather than the standard method — which for
copper (II) hydroxide is chemical simplicity.
Byda - what is your electrolyte?Squall - 12-6-2010 at 15:49
My electrolyte consists of distilled water and NaCl, i chose this method because it seemed to be a simple way of making copper hydroxide, my actual
goal is to make CuO. The WiZard is In - 12-6-2010 at 16:57
By "you people" you mean those of us less worthy than your esteemed self? What a great entertainment you
are! We love you old fahrts.a_bab - 13-6-2010 at 00:37
If copper oxide is what Squall is after and he uses NaCl in the eletrolyte it means he is not concerned at all about possible Na contamination (thus
he doesn't need the CuO for pyro use I'd say).
Having that, why he doesn't just use the "first grade studied" reaction Wiz is proposing and instead farts around with electrolysis it's beyond my
comprehension. Some people just like fly over to cook an egg into a volcano crater for the breakfast it seems.
While I enjoy entropy51's posts and he's knowledge I guess he is in no position of saying anything about Wiz. If I were to say who's the spammer on
this thread, it's entropy51. No contribution at all here.
[Edited on 13-6-2010 by a_bab]Squall - 13-6-2010 at 04:53
I know one could go and mix sodium hydroxide with copper sulfate and be done with it, but why not explore the other possibility of electrolysis. I
know that this is probably not the best or even a cost effective method but my reasons for performing this is more for the knowledge gained then for
the end product. As I am performing these trials I am learning about different electrolytes and their products, and I would like to thank you for your
advice and criticism it has been helpful in pointing me in the right direction. Maybe someday I'll Learn how to post like the pro's but that's in the
future.
Anyway I have switched electrolytes, because several of you made the comment about the presence of chloride in the electrolyte, I will now use sodium
bicarbonate and see what happens. Thanks again12AX7 - 13-6-2010 at 05:03
Hah, copper tends to form a carbonate complex. You will eventually have a deeply blue solution (and some copper plating over, and not much copper
carbonate!). This is also a poor way to generate copper hydroxide.
Your best bet is an anion which doesn't reduce easily and doesn't complex with copper significantly. Nitrate and chlorate do, so they are out.
Acetate forms a complex, so it's out. Perchlorate and most accessible of all, sulfate, come to mind.
TimSquall - 13-6-2010 at 05:12
Sulfate is accessible, but what I am not sure of is how much sulfate contamination will be present in the hydroxide, if it is minimal then that's the
way I should go.
Thanks Timnot_important - 13-6-2010 at 05:36
Might try the following:
copper electrodes
low voltage AC - not DC
whatever electrolyte
brisk stirring
running hot, near 90 C
Let's say you use NaCl. At the (momentary) anode Cu(I) and Cu(II) are formed. At the (momentary) cathode H2 and OH(-) are formed. The mixing insures
the copper ions and hydroxide meet, forming copper hydroxides. On reversal of polarity some copper ions will plate out onto the cathode, but will be
stripped off on the following polarity reversal. With NaCl some hypochlorite and chlorate are formed, the OCl isn't stable under the conditions both
oxidising Cu(I) and disproportionating to chloride and chlorate. The hot conditions encourage Cu(OH)2 to convert to CuO, but there's still be a mix
in the precipitate; that can be taken care of later.
If you use NaHCO3 electrolyte, you'll get a mixture of copper hydroxides and basic carbonates. Ammonium sulfate will give mostly CuO with some basic
salts, the electrolyte will lose some ammonia so you may have to add aqueous ammonia to keep the pH from going too acid. Yes, some copper will remain
in solution, how much depends on the electrolyte salt used, but it's kind of a 'who cares'; you can save and reuse the electrolyte if you want.
In all of those cases, after you stop the electrolysis let the rig cool and precipitate settle. Decant the electrolyte off, wash with a little water,
add a fair amount of water then boil for some minutes. Cool and settle, decant through a filter, wash several times by decantation - again pouring off
the water through the filter to catch bits of escaping oxide. Finally wash the rest of the oxide into the filter and give a final rinse. Let the oxide
air dry, crushing big lumps as it does so. Put the oxide in a glass or ceramic dish or bowl style container, evaporating dish if you have one of
proper size. Slowly heat it with stirring, and lump crushing, to at least 120 C for non-carbonate electrolytes, for carbonates and preferably for
others to at least 300 C to insure full decomposition to CuO.
Note that technical CuSO4, such as 'root killer', may contain appreciable amounts of iron. If you wish to go the CuSO4 + NaOH route to the hydroxide,
you best first test for iron and if found purify the CuSO4; you can't do this by recrystallising, chemical means must be used. The WiZard is In - 13-6-2010 at 06:08
Sulfate is accessible, but what I am not sure of is how much sulfate contamination will be present in the hydroxide, if it is minimal then that's the
way I should go.
Thanks Tim
---------
Try ammonia water (ammonium hydroxide).
Active Nature of Copper Proved in Your Laboratory
Popular Science
June. 1934
If i was to pursue the copper sulfate + a hydroxide , my only source at the moment of copper sulfate is root killer, my question is how do I test it
for iron content and if it is contaminated by what chemical processes do I remove it.
And thanks for the links Tim.not_important - 13-6-2010 at 07:28
One way -
Make a fairly strong solution of a bit of the root killer, just a crystal or two in 1 or 2 ml water. Add several ml of hydrogen peroxide %3 or a few
drops of a higher concentration H2O2. Put a drop or two of this on a piece of filter paper, or even plain paper toweling. Drop concentrated aqueous
ammonia onto that drop, slowly adding several more drops right on top. Copper hydroxide forms at first, then the excess ammonia complexes with it and
dissolves the copper away, forming an expanding circle of blue becoming paler and paler as additional ammonia water dilutes it. Iron will form the
hydrated oxide, and remain as a rusty spot where the drop of root killer solution was originally placed.
The clean-up is related. Make a solution of root killer, add H2O2 and heat to near boiling for 10 to 15 minutes, adding a few ml of H2O2 every minute
or so; the set aside to cool. When it's lukewarm measure out about 10% into a second container.
Take that 1/10 and add aqueous ammonia to get Cu(OH)2, slowing down the addition as the solution appears to get colourless (let a couple of drops drip
through filter paper if you need to. It's not real critical to get it exactly right on balanced, a bit to little NH3 or a bit over isn't going to
hurt, but not a whole lot too little or much. Filter and wash the Cu(OH)2 with distilled water.
Add the washed Cu(OH) to the main solution of root killer, and stir it for awhile. Fe(III) hydroxide/hydrated oxide is very insoluble, Fe(III) salts
in solution will exchange with the Cu(OH)2 to give a precipitate of "ferric hydroxide" with the copper going into solution as whatever the iron salt
was - sulfate in this case.
After stirring for at least 15 minutes to a half hour, you can take let it settle for awhile and then test a drop of the solution - free from
precipitates - for iron as before. If it still tests positive, repeat the H2O2-boil steps, but only use about 5% for making the Cu(OH)2.
Filter the solution if the precipitate looks to be easy to filter. If the ppt is sort of floaty then try bringing the solution to a boil again to
convert the excess Cu(OH)2 to the oxide, which generally is easier to filter. The iron is mixed in with the excess copper hydroxide/oxide.
You can just let the filtrate, covered with a cloth to keep out dust, evaporate slowly until it's mostly crystals with a cm or so of solution above
it. Or you can heat it to near boil, until crystals start to form, add enough hot distilled water to dissolve those then let the hot solution get
cold. This will give you about 3/4 of the copper sulfate as crystals, after filtering off the mother liquor you can take that and concentrate it to
et a second batch of crystals. Don't evaporate to dryness, expect to leave 5% to 10% of the CuSO4 in solution along with whatever impurities; iron
sulfate would have formed mixed crystals so that's why it needed to be removed chemically.
If all you want the CuSO4 for is making CuO, you can skip crystallisation and just used the filtered solution as is.
[Edited on 13-6-2010 by not_important]The WiZard is In - 13-6-2010 at 09:11
Hah, copper tends to form a carbonate complex. You will eventually have a deeply blue solution (and some copper plating over, and not much copper
carbonate!). This is also a poor way to generate copper hydroxide.
Tim
--------- Poor way? Actually it is THE way. Brauer add's a detail to the
process - ammonia water is added first.
La Book is in the SicMad library.Hamilton - 13-6-2010 at 11:44
i use to make pound of CuO straight by electrolysing Cu metal in pool calcium hypochlorite or simple bleach. The content of cuO over Cu2O was about
40:1 in weight. another way is to electrolyse chlorate solution with cu metal, but then the ratio of CuO to Cu2O is 15:1 or even worst.
You can easily separate the two with some acetic acid as only CuO dissolve. Then if you still wan CuO instead of CuAc then you are screwed ;-). but
CuAc is like CuO if you intend to reac the CuO with strong acid like HCl of H2SO4...
it seem you can convert Cu(I) to Cu(II) by dissolving it in excess HCl but i don't know why, it just look like because it create a green solution
ressembling CuCl2.
used to have a blog about my copper experiments but it got corrupted when the host broke his computer
[Edited on 13-6-2010 by Hamilton]
[Edited on 13-6-2010 by Hamilton]woelen - 13-6-2010 at 13:07
I can imagine that making copper hydroxide in this way is an interesting option. Copper wire is something which is available everywhere. Copper
sulfate more and more becomes a less common chemical, although of course it still is available through eBay and other online sources.
A nice method of making Cu(OH)2 is electrolysing a solution of baking soda (NaHCO3) with copper electrodes. Quite some precipitate of light blue
Cu(OH)2 is formed. This precipitate also may contain some CuCO3. If you filter the precipitate and heat it, then it becomes black and CuO is formed.
Even the wet suspension in water can be boiled to make a black precipitate of CuO. Any carbonate then is destroyed and expelled as CO2. The CuO,
prepared in this way is not suitable for pyrotechnical purposes, due to sodium-remains (the orange/yellow light if sodium overwhelms the cyan color of
copper in pyrotechnic flame compositions). For many other chemical experiments (e.g. making CuCl2 by dissolving it in dilute HCl) it is perfectly
suitable.
[Edited on 13-6-10 by woelen]bbartlog - 13-6-2010 at 15:38
'it seem you can convert Cu(I) to Cu(II) by dissolving it in excess HCl but i don't know why'
I don't think this works. HCl alone will not oxidize CuCl to CuCl2 (though it will dissolve it by forming complexes). Now, in the presence of air,
oxygen will gradually oxidize Cu(I) to Cu(II), which will immediately be converted to CuCl2.
I've made Cu(OH)2 by first dissolving copper in HCl to make CuCl2 (air has to be bubbled through for this to work), then reacting the CuCl2 solution
with NaOH. CuCO3 is also easily produced via similar neutralization with Na2CO3.
So far as the yellowish substance mentioned in the first post is concerned, it sounds like cuprous hydroxide (see: http://pubs.acs.org/doi/abs/10.1021/j150103a002). For what it's worth, cuprous carbonate (or whatever you get when neutralizing a solution of CuCl
and HCl with Na2CO3) is an orange insoluble precipitate. Exposure to air turns it green.
