[Xeluc]

Ok, I set up and experiment where I disolved NaCl into H2O until it was saturated. I then got a 9 Volt DC power supply and hooked up a Copper Anode and a Graphite Cathode. When the power is turned on, alot of gas bubble from the Graphite Cathode. I am assuming that this is H2. Relatively fewer bubbles are on the cathode. (Like hardly anything bubbling up at all.) The purpose of this experiment was to Create a Copper Salt. I was looking for CuCl2. Anyhow this is what happened. An orange precipitate came from the Anode. (arent most copper compounds green or blue?) Later in the reaction I noticed some seperation. The precipitate had floated on top of some water. By this time the water was a faint green color (obviously some sort of copper Ion or compound) I havnt isolated the green compound yet, but I am more focused on this orange precipitate. I conducted a flame test which told me that it was Sodium based. It isnt NaOH becasue that is white. I boiled off the water and was left with a dirty yet pure orange powder. I have no idea what this is. In a few days i will switch between graphite and copper cathodes and anodes to see what combinations will produce this precipitate. Until then, can anyone help try to figure out what this compound is? Thanks for any help I get, I know this was long

[Mitch]

It's sodium hydroxide with copper residues in it.

[Xeluc]

im not sure that its NaOH because this substance is insoluable in water... it settles.. If im not realizing something though please correct me.

[woelen]

Ok, I set up and experiment where I disolved NaCl into H2O until it was saturated. I then got a 9 Volt DC power supply and hooked up a Copper Anode and a Graphite Cathode. When the power is turned on, alot of gas bubble from the Graphite Cathode. I am assuming that this is H2. Relatively fewer bubbles are on the cathode. (Like hardly anything bubbling up at all.) The purpose of this experiment was to Create a Copper Salt. I was looking for CuCl2. Anyhow this is what happened. An orange precipitate came from the Anode. (arent most copper compounds green or blue?) Later in the reaction I noticed some seperation. The precipitate had floated on top of some water. By this time the water was a faint green color (obviously some sort of copper Ion or compound) I havnt isolated the green compound yet, but I am more focused on this orange precipitate. I conducted a flame test which told me that it was Sodium based. It isnt NaOH becasue that is white. I boiled off the water and was left with a dirty yet pure orange powder. I have no idea what this is. In a few days i will switch between graphite and copper cathodes and anodes to see what combinations will produce this precipitate. Until then, can anyone help try to figure out what this compound is? Thanks for any help I get, I know this was long

Copper chemistry is very complicated, especially, when both copper (I) and copper (II) chemistry are involved. At the link below, I give a description of a set of experiments I recently did with copper. These really puzzled me, but now I have quite some idea what is happening.

The yellow/orange stuff you get is a copper (I) compound. Under the given conditions, apparently the copper wire is not oxidized to copper (II), but to a copper (I) compound. More likely it is a copper (I) compound, contaminated with some copper (II).

I think that at the anode you have the following reaction:

Cu + OH(-) --> CuOH + e(-)         (OH(-) from cathode reaction)
2CuOH --> Cu2O + H2O

In fact the two reaction can even better be described as formation of hydrous copper (I) oxide with composition Cu2O.xH2O. Cu2O.xH2O is orange/yellow. When contaminated with copper (II) it is dirty brown/yellow. Anhydrous Cu2O is brick red.

What I do not (yet fully) understand is the role of the chloride in the electrolysis experiment. When you perform electrolysis of a solution of NaOH, then you get O2 at a copper anode! When, however, quite some NaCl is mixed in, then you get the yellow copper (I) compound. Apparently, the Cl(-) plays an important role by means of complex formation, changing the redox-properties of copper considerably. In the presence of Cl(-), the oxidation state +1 of copper is strongly favoured (see links on webpage below).

This hydrous copper (I) compound can also be prepared chemically, without electrolysis.

Look at the pictures of the end of the webpage. Probably this brown/orange/yellow compound is very similar to what you have made by electrolysis.

    http://www.woelen.nl/chem/exp0005/exp0005.htm

The yellow stuff is a copper (I) compound. When some concentrated HCl is added to the yellow stuff, then a dark solution is obtained, which becomes bright green on addition of some H2O2. The H2O2 oxidizes all copper (I) to copper (II), which in turn forms a green complex with chloride.


Wilco

[Xeluc]

Well, my experiemtns are definatly coinciding with yours. The precipitate looks almost exactly like mine. thing is.... it doesnt stay there if the electrolysis is continued. a dark liquid eventually takes over.. jsut like in your experiment. I  have realized that there are so many reactions taking place that this isnt accomplishing my main reason of doing this. I need conditions that restrict reactions. I need one of those semi-permeable membranes to seperate the solutions in. thing is... NaOH for example could still be made in either solution.. i think... NaCl wold split up and the Na would head over to the cathode... well if a hydroxide ion bonded to it b4 it crossed the ion permeable meambrane.. thaen itd be on the anode side.. if it got through first.. then itd be on the cathode... so maybe that isnt a great idea. thing is.. I found a better way of doing what i wanted to in the frist place (create CuCl2) I figure if i do electrolysis in NaCl water using graphite cathode and anode.. i shoul be able to get Cl2 to accumulate at the anode. it wouldnt be very hard getting it into a contaner and adding some copper... Would that be an explosive reaction? I was thinking so seeing as copper isnt very reactive as aposed to Na which makes quite a show with Cl. So i guess im going to make an electrolysis thing that can channel Cl2. will Cl2 melt through plastic? i hope not...

[Grafter]

Cl₂ won't corrode rubber tubing, but i'd reccomend a glass set for it anyhow. Cl₂ is toxic, and I wouldn't reccomend it for non-lab use, and without a proper fume hood.

Secondly, obtaining Cl₂ gas from an aqueous solution may be problematic. The oxidation potential of Cl^- is -1.358v. At pH 7, the oxidation potential of H₂) is about -0.8. This means that to put in the potential that you need to oxidise Cl^-, you will oxidise H₂O to oxygen first. Try doing it in a solvent that is more resistant to oxidation.

Grafter

[jdurg]

The thing is, hydrogen is produced at the anode and chlorine at the cathode.  So you need to see what the potential is for oxygen to form rather than hydrogen.  The standard oxidation potential for water is -1.23 volts while that for Cl- is -1.36 volts.  That is so close that in reality, the chloride is more easily oxidized than the water is so chlorine gas forms at the anode and very little oxygen is formed (Due to the overvoltage).  So what happens is that you get a solution of sodium hydroxide from the reduction of water at the anode, and chlorine gas from the oxidation of chloride at the cathode.  Some Cl2 will probably come up out of the solution, but a lot of it will redissolve into the NaOH solution giving you sodium hypochlorite.  Basically, you wind up with bleach.  

[Xeluc]

The thing is, hydrogen is produced at the anode and chlorine at the cathode.  So you need to see what the potential is for oxygen to form rather than hydrogen.  The standard oxidation potential for water is -1.23 volts while that for Cl- is -1.36 volts.  That is so close that in reality, the chloride is more easily oxidized than the water is so chlorine gas forms at the anode and very little oxygen is formed (Due to the overvoltage).  So what happens is that you get a solution of sodium hydroxide from the reduction of water at the anode, and chlorine gas from the oxidation of chloride at the cathode.  Some Cl2 will probably come up out of the solution, but a lot of it will redissolve into the NaOH solution giving you sodium hypochlorite.  Basically, you wind up with bleach.  

EXACTLY! the only problem is containing it.. chlorine sinks in air if im correct.. am i correct?
so i dont know how ill do this
i need to brain storm how to get chlorine up and over w/out it jsut backing back out the tube. since pressures isnt being relieved when the Cl2 is made, it will jsut back out.. so anyone know how to go about this?

[woelen]

EXACTLY! the only problem is containing it.. chlorine sinks in air if im correct.. am i correct?
so i dont know how ill do this
i need to brain storm how to get chlorine up and over w/out it jsut backing back out the tube. since pressures isnt being relieved when the Cl2 is made, it will jsut back out.. so anyone know how to go about this?

If you want to make CuCl2, then you'll probably have a hard time to get it in a reasonably pure state. Coper metal reacts with chlorine, non-violent if copper wire is added to moist chlorine. I would, however, not go through the chlorine gas. It is quite dangerous and very tedious if you need electrolysis to prepare some of the gas.

Another, much easier way is to add copper wire to concentrated hydrochloric acid and add in some 30% H2O2, very slowly. The copper first dissolves quickly and the liquid becomes green. Lateron, the chlorine dissolves more slowly and the liquid becomes dark brown. If the latter occurs, then drop in H2O2, until the liquid just has become green again. In this way, you can make very concentrated solutions of copper chloride in HCl. By heating the liquid, you can drive off most HCl and water, but it will be hard to get rid of the last traces of HCl. The problem is that CuCl2.2H2O does not crystallize easily. It is VERY soluble and quite hygroscopic.

This preparation MUST be done outside. You get a lot of noxious fumes.

If you do not have 30% H2O2, then you could use bleach instead, but in that case, you need to be even more careful. VERY slowly, with constant swirling add bleach to the liquid. The swirling is really important, otherwise you will get a lot of bubbles of chlorine gas. On the other hand, the latter method of preparation probably is too dangerous to be attempted without good lab equipment. There is the risk of evolution of a lot of chlorine gas and beware, chlorine gas has killed many people before!

Yet another and MUCH safer method is the following. Go to a ceramics and pottery supplier and get some copper carbonate (nice cyan/green powder) or copper (II) oxide (black powder). Add this to luke warm hydrochloric acid, until all has dissolved (the copper carbonate gives bubbles of carbon dioxide). You can dissolve quite some copper carbonate or copper oxide in a small volume of concentrated hydrochloric acid. The resulting liquid must be evaporated to drive off the remaining HCl and water, but again, it will be hard to get a nice dry product. Heating too strongly will not help you, because that will result in partial decomposition and you end up with copper oxychloride, a basic copper chloride, which is insoluble in water. If you drive off HCl, do that outside. The fumes are quite corrosive and if you do that inside, then everything made of iron in the neighborhoud will rust away.

Wilco

[Xeluc]

All I can say is wow.. Thank you so much... i LEARNED about carbonates in acid just this year.. stupid stupid stupid.... man I am learning so much on this forum, it's great. As I dont really want to order any more chemicals online I'd like to go with either of the first two. I have bleach, but not 30% H₂O₂. Is there a way or concentrating H₂O₂? Im going to google this right now. thank you so much for your help, that was everything I needed and more. The awesome thing is, this information can be used for alot of things other than CuCl₂. I just have this thing where I think its really neat to manufacture something out of other things, especially simple ionic compounds. So thats that. Now back to my orange precipitate. you told me that it is a Cu (I) compound.. Well Adding H₂O₂ and HCl to it should yeild CuCl₂ correct?

[woelen]

All I can say is wow.. Thank you so much... i LEARNED about carbonates in acid just this year.. stupid stupid stupid.... man I am learning so much on this forum, it's great.

I'm quite sure that you will learn more in the future. There are many enthusiastic and nice people on this forum...

As I dont really want to order any more chemicals online I'd like to go with either of the first two.

Can't you get it locally? Pottery and ceramics shops are quite abundant, at least where I live.

I have bleach, but not 30% H₂O₂. Is there a way or concentrating H₂O₂? Im going to google this right now.

I read about freezing out the water. If you put 3% H2O2 in a freezer and let is freeze partially (such that approximately half the liquid is frozen), then you have a higher concentration in the non-frozen part. By repeating this process twice it should be possible to get 10% H2O2. That should be OK for most of your experiments. Be careful though with H2O2 of more than 5% concentration. It destroys your skin almost instantly. it gives white spots, which sting like hell. I've personal experience with that!

thank you so much for your help, that was everything I needed and more. The awesome thing is, this information can be used for alot of things other than CuCl₂. I just have this thing where I think its really neat to manufacture something out of other things, especially simple ionic compounds. So thats that.

Now back to my orange precipitate. you told me that it is a Cu (I) compound.. Well Adding H₂O₂ and HCl to it should yeild CuCl₂ correct?

Yes, first add HCl, such that it dissolves. You get a dark green/brown liquid. Next add H2O2 dropwise until the liquid is nice bright green. For this, you can use 3% H2O2 and 10% HCl. It might be that even after adding the HCl, it is green already, because oxygen from the air also oxidizes the copper (I) compounds easily.

Wilco

[Xeluc]

ok great, as soon as im home ill try this. of course like everything else ive thoguht would work, nothing has. heres why I think this. I conducted a flame test on that compound. had it been a copper compound it sohuld have flamed green. It flamed orange/yellow, which is what sodium does. So.. unless yet again my logic is flawed here, this cant be a copper compound.

another thing that caught my attention is this. using a graphite anode in the reaction yeilded Cl₂. This is cuz the Cl₂ has nothing to reacti with so it bubbled up, possibly with O₂, but that is irrelevant. I knew what it was cuz I could smell it. When it is replaced by a copper anode, i do not smell Cl₂. It would be reasonable to say that the copper reacted with the chlorine then, correct? if so, that I HAVE to making a copper chloride-like compound. I guess it could be hypochlorite also, but

[Borek]

Sodium traces will mask copper. You started with brime and I doubt you will be able to remove sodium from the substance in question.

[Xeluc]

ok... i went up and bought some H₂O₂. ok... I put hcl into my orange powder and it turned green. the addition of H₂O₂ did nothing but dilute the solution. as a control, i also mixed H₂O₂ and HCl together and added copper wire. It also turned green. So, looks like oyu guys were right. it took a long time, but i think i figured out how to synthesize CuCl₂. Thanks alot guys.... I agree that the Na will be hard to get out of the 1st solution. I dont plan on using that for CuCl₂. I was jsut seeing if it was a copper compound. The 2nd batch I made should be reasonably pure, and i intend on boiling away the water/HCl. I only have a small butane torch to do it though.... Anyhow, here is my last question: How does the addition of H₂O₂ disolve the Copper. The way I see it, If copper won't disolve in HCl because it isn't reactive enough, then by rights the same thing should happen with H₂O₂. In fact this is true. Not one of them alone will have an effect, but together they disolve the Cu into CuCl₂. Could someone please explain why this is so? thanks

[Borek]

H⁺ ions present in HCl solution are too weak (they have too low electronegativity) oxidizers to dissolve copper. H₂O₂ is strong oxidizer, especially in acididc solutions. That's why copper dissolves after H₂O₂ is added.