More copper electrolysis fun!

[jsatan]

So Now I've got a better power suplly I didnt the age old thing.

Salty water and copper electrode.

 

You know the story,

You get (blue) copper hydroxide and (brown) copper oxide.

The copper hydroxide will decomposses if over a certain temperature.

 

So I thought lets cool it and try and make just copper hydroxide.

So I Let the salty water sit in the freezer until nice and cool and then placed in a ice bath just to make sure.

 

But I didnt get copper hydroxide.

I got a green/brown precipitate.

 

I'm sure I made copper chloride.

I tested this by taking a sample of the sludge and adding Sodium hydroxide.

I got copper oxide precipitate and copper hydroxide soln

I did the same with my sample of copper oxide as a control test, No reaction.

I also notice if I dont descent a sample it will (as it warms) turn into copper hydroxide(which is should anyway).

 

There is one thing I did notice with all this, is that the salty water and sodium hydroxide will form different layers.

This can be used to change what is made.

 

If I've missed any information out let me know as I've had to type this 2x as the library computers are.....(fill in blank).

[xeluc]

Ok, lets disect this message and see what we can do. I love Copper, you can do so much stuff with it and it's really cheap. Anhow...

 

 

You know the story' date='

You get (blue) copper hydroxide and (brown) copper oxide.

The copper hydroxide will decomposses if over a certain temperature.

[/quote']

 

For SOME reason, I get only Cu2O, maybe Woelen can save the day and explain our differing outcomes. I use a high potential (12 Volts) so maybe that makes a difference. Anyhow, this is irrelevant. You are correct, Copper Hydroxide will dehydrate to form CuO. But watch, here it looks like you understand that the Copper Hydroxide is a precipitate and therefor insoluble. Keep this in mind, we will go back to it later.

 

So I thought lets cool it and try and make just copper hydroxide.

So I Let the salty water sit in the freezer until nice and cool and then placed in a ice bath just to make sure.

 

 

Although you are correct in saying that heat will decompose the Cu(OH)2' date=' The Copper oxide formed during electrolysis may or may not be from heat. Voltage may also contribute. This is just something to think about, I go in work in an hour and can't test this, but if noone has by tomorrow I will.

But I didnt get copper hydroxide.

I got a green/brown precipitate.

 

No idea here. Are you sure only copper metal is touching the solution?

 

I'm sure I made copper chloride.

I tested this by taking a sample of the sludge and adding Sodium hydroxide.

I got copper oxide precipitate and copper hydroxide soln

 

WOAH' date=' stop right there. remember that thought I told you to hold onto on the first quote? Well right here, you forget that Copper Hydroxide is insoluble in water. It is possible however that you had some Cu2+ ions in solution which would crystalize CuCl2. I have done this myself actually. So after your test, you should have blue Copper hydroxide and copper Oxide maybe. If you SOLUTION is blue, then You have something else. Woelen told me that you would get a royal blue solution of the CuO2(2-) complex.

I did the same with my sample of copper oxide as a control test, No reaction.

I also notice if I dont descent a sample it will (as it warms) turn into copper hydroxide(which is should anyway).

 

The first part of this is absolutly right. Next you say that warming your mixture turns it into Copper Hydroxide? This obviously can't be true since you said already that heating copper hydroxide yeilds Copper oxide. This isn't a reversable reaction with heat or lack of alone.

 

 

 

 

So anyway, Electro chemistry is a finicky thing. In truth, when you first start out, you spend more time figuring what went wrong than conducting something that you can successfully predict. I hope I've shed some light on this, but I don't have a lot of time to experiemtn right now. So, I'll do that tonight or tomorrow. Hope I helped a little

[woelen]

JSatan, as Xeluc already pointed out, he gets Cu2O instead of Cu(OH)2. From your observation, I can only conclude that you also get Cu2O, probably contaminated with some copper (II) compound.

You write that you get a green/brown precipitate at the anode. This brown color can be explained by assuming that you get orange/yellow copper (I) oxide, heavily contaminated with a copper (II) compound. In this complex mix you'll also have multi-valency compounds, which are dark brown. Even tiny amounts of these compounds in your precipitate strongly affect its color.

 

By means of electrolysis of a solution of NaCl with copper electrodes it is amazingly difficult to get a nice copper (II) compound, you get all kinds of complexes and this is due to the presence of chloride in high concentration. Chloride ion is a very strong coordinating/complexing agent and together with the fact that copper (I) is stabilized strongly by chloride-ligands it can be easily understood why the formation of copper (I) compounds is favored over formation of copper (II) compounds.

 

If you want to make nice copper (II) compounds by means of electrolysis, then I would suggest you to use a slightly acidified solution of Na2SO4 (acidified with H2SO4 or HNO3, NOT with HCl ).

 

Pure copper hydroxide is really blue, like the blue sky on a bright and sunny winterday. Any brownish or greenish hues tell you that there are impurities in your product.

Copper hydroxide will turn into almost black CuO on heating, even when immersed in water. On cooling down, it will not revert to Cu(OH)2.

[jsatan]

aha, I'm getting there slowly but its fun on its way,

Yeah the blue/white precipitate with room temp water did make me think.

I was using 12V (I know its too much but its all I've got atm).

Now I know What I'm getting that will help.

I think I remember That there is something good about copper hydroxide but I'm not sure what it was.

 

I'll be back with more later in the week,

Even tho it didnt do anything that I thought I've still learned something, and thats the best bit,

[jsatan]

The first part of this is absolutly right. Next you say that warming your mixture turns it into Copper Hydroxide? This obviously can't be true since you said already that heating copper hydroxide yeilds Copper oxide. This isn't a reversable reaction with heat or lack of alone.

lol. Nah I dont mean heating I mean letting it warm to room temp,

Not that it matters now,

Yeah I had the copper probes in the water all the time.

I did get to a point with one sample of black crap plating out, lol.

Thats whne I knew the soln was full.

[xeluc]

lol. Nah I dont mean heating I mean letting it warm to room temp' date='

Not that it matters now,

Yeah I had the copper probes in the water all the time.

I did get to a point with one sample of black crap plating out, lol.

Thats whne I knew the soln was full. [/quote']

 

Even letting it go to room temp wont change anything to Copper Hydroxide...

 

 

@Woelen: Actually, it's not easily undertood by me at least why Copper (I) Compounds ar/ efavored over Copper (II) in a NaCl solution. I understand that the Cu+ is stabailized by teh Chlorine Ligands but Cu2+ (obviously) does the same. Does voltage have anything to do with the preference? Maybe you could elaborate on this.. Thanks

[DukerX]

What does the (I), (II) and (III) mean when you talk about CuO?

I noticed this remark about making "nice copper (II) compounds"

How "nice" would these be, and could it be done whit KNO3 instead of Na2SO4, as I have no idea of where to even start looking for Na2SO4, but I have LOADS of KNO3.

 

Another thing that puzzles me is how to separate the CuO from the Na2SO4 or perhaps KNO3 when I'm "done"...

I was searching for "Na2SO4" in my native language and found some crude scool-reports on electrolysis of Na2SO4 using carbon electrodes.

It was said that after a while, the solution got a higher pH (alkaline??).

Would I have to counter this whit more H2SO4, or would the increase in pH be insignificant?

 

It would be great to be able to make alot (hundreds of grams) of uncontaminated "Copper(II) Oxide" as I need it in some of my other "chemistry" experiments.

[akcapr]

The numeral indicates the charge on a variable-cahrge metal. Metals such as Iron can have variable charges- II (+2) and III (+3) for example.

 

You can seperate by filtering or sometimes crystillization. ANd the increase in pH would screw up your electrolysis how? so it shouldnt matter. However depens wat the purpose of the electrolysis is.

[woelen]

What does the (I)' date=' (II) and (III) mean when you talk about CuO?

I noticed this remark about making "nice copper (II) compounds"

How "nice" would these be, and could it be done whit KNO3 instead of Na2SO4, as I have no idea of where to even start looking for Na2SO4, but I have LOADS of KNO3.

[/quote']

With "nice" I meant pure. With NaCl as electrolyse it will be very difficult to obtain even moderately pure compounds. Having a mix of all different kinds of compounds which you do not know cannot be regarded as "nice".

With KNO3 you also can do the trick, although impurities formed at the cathode probably will spoil things somewhat. Electrolysis of KNO3 results in formation of H2 and OH(-) at the cathode, but as a side reaction, you will also get NH3 at the cathode. The nitrate ion can be reduced to NH3 under strongly reducing conditions and that is exactly what you have at the cathode. The NH3 also forms a complex (deep blue tetramine copper (II)) and that again may result in unwanted contamination.

 

Another thing that puzzles me is how to separate the CuO from the Na2SO4 or perhaps KNO3 when I'm "done"...

I was searching for "Na2SO4" in my native language and found some crude scool-reports on electrolysis of Na2SO4 using carbon electrodes.

It was said that after a while, the solution got a higher pH (alkaline??).

Would I have to counter this whit more H2SO4, or would the increase in pH be insignificant?

In your situation with lots of KNO3 at hand, forget about the Na2SO4. I mentioned this as just being a rather inert salt, which does not form complexes with copper (II) ions, nor with copper (I) ions. For most people, Na2SO4 is easier to obtain than KNO3, so I mentioned that.

Separating CuO from the liquid is easy. CuO forms a precipitate and sinks to the bottom. Decant the solution, add a lot of water, let precipitate settle again and decant the water again. Doing this twice gives you wet CuO, which can be dried in a few days on a piece of filter paper, which you put in a warm and dry place.

 

It would be great to be able to make alot (hundreds of grams) of uncontaminated "Copper(II) Oxide" as I need it in some of my other "chemistry" experiments.

A very nice source of chemicals are pottery and ceramics suppliers. They have many metal oxides, metal carbonates and metal sulfates at reasonable purity and really cheap. If I were you, I would try to grab some metal salts from such a shop.

 

 

With copper (I) I mean compounds of Cu(+), e.g. CuCl and Cu2O. With copper (II) I mean compounds of Cu(2+), e.g. CuCl2, CuSO4, and CuO.

[DukerX]

I live in this rockpile of a country called Norway.

Here, everything out of the ordinary toothpaste, milk and sugar is surprisingly hard to get hold of, unless you have "connections".

If I'm at all able to find a local pottery supplier, it's doubtfull anything they have would come across as "cheap".

Anyhow... I'll look into it. Maybe I'm up for a surprise...

[xeluc]

If you get two cups, fill each with KNO3 solution, get some yarn (maybe 3-4 strings wide) and wet it in the solution and drape it between the cups so that the liquid make a "connection" between the cups. This is basicly a Salt bridge. What this does is keeps contaminates form the Cathode (In this case, the Tetramine Copper complex solution) from getting into your Anode solution, which will be your CuO (I Think. Woelen said the Cl-'s from the NaCl solution favor Cu+ Ions, so with KNO3, you should get CuO?)

[woelen]

If you get two cups, fill each with KNO3 solution, get some yarn (maybe 3-4 strings wide) and wet it in the solution and drape it between the cups so that the liquid make a "connection" between the cups. This is basicly a Salt bridge. What this does is keeps contaminates form the Cathode (In this case, the Tetramine Copper complex solution) from getting into your Anode solution, which will be your CuO (I Think. Woelen said the Cl-'s from the NaCl solution favor Cu+ Ions, so with KNO3, you should get CuO?)

No, you will not get CuO (or Cu(OH)2) in this way. You actually need the cathode material. At he anode you get Cu(2+) ions and with OH(-) ions, formed at the cathode, you get your Cu(OH)2.

 

Anode: Cu --> Cu(2+) + 2e (absorbed by power supply)

Cathode: 2H2O + 2e (given by power supply) --> H2 + 2OH(-)

 

When the liquid is mixed (and it does on its own, due to the motion, induced by the rising gas bubbles), then the hydroxide and copper ions form a precipitate.

[jowrose]

I tried electrolysing copper, and all I got was a big mass of brown sludge. Is it copper oxide then? My whole goal is to make copper chloride, so I don't really care what compound I get as long as it will react with hydrochloric acid. I mixed this unknown gunk with HCl, and lo and behold I get an almost pumpkin-colored solution. Not quite the color I want... What am I doing wrong?

[xeluc]

When you have a mix of Cu+ and Cu2+ in HCl it makes a brown solution.. I can tell you how to synthesize CuCl2 though if by chance you have made something else. Put Cu in HCl. Add some H2O2. You'll notice the solution turns green. Leave the Copper in longer, it turns brown. No worries, its just Cu+. It will oxidize to Cu2+ in air. If you then just let the whoel thing sit around for a few days. The CuCl4 2- in solution will eat away more Copper, resulting in CuCl2 - which ozidizes in the air, only to react with more copper. So it's a continuous reaction. When your finiahed, either add more H2O2 to make the solution toatlly green, or let the solution set out in air without copper in it and it will slowly turn green. Theres your CuCl2. As long as you used a salt bridge of some sort, I don't see how having HCl and a copper Anode wouldn't also create your CuCl2, but I've never done it. I have Electrolized Cu in NaCl solution, and I get precipitates? HOW?

[jdurg]

I have Electrolized Cu in NaCl solution, and I get precipitates? HOW?

 

The electrolysis of an NaCl solution produces chlorine gas and sodium hydroxide. As the copper metal corrodes away, the copper ions will react with the free OH- ions forming Cu(OH)2 which will ppt out of solution.

[jowrose]

mmmm chlorine.

 

Ok, I'm a little confused about the whole electrolysis reaction. You can electrolize water, and you get hydrogen and oxygen (add a spark for fun). But when you add NaCl to the solution, do you get H2, O2, and Cl2 gases? and then the sodium by itself?

[xeluc]

The electrolysis of an NaCl solution produces chlorine gas and sodium hydroxide. As the copper metal corrodes away, the copper ions will react with the free OH- ions forming Cu(OH)2 which will ppt out of solution.

 

Yeah I know that... I was using a salt bridge. No OH's present.. so now how? lol

 

@Jowrose: under normal circumstances... H2 and Cl2 are produced I believe. O2 will be produved at higher voltages maybe

[jowrose]

Why is the O2 not produced? Does it immediately oxidize the electrode? Because the H2 can't come from nowhere...

[jdurg]

Yeah I know that... I was using a salt bridge. No OH's present.. so now how? lol

 

@Jowrose: under normal circumstances... H2 and Cl2 are produced I believe. O2 will be produved at higher voltages maybe

 

 

Yes, but in the chamber with the copper hydrogen gas would form. That would leave you with Na+ and OH- ions in solution. Copper would then go into solution to balance the charges. No matter how hard you try, OH- will exist in solution. Otherwise it would be impossible for the Hydrogen gas to form.

 

@jowrose: From what I understand, xeluc was using two separate container connected with a salt bridge. The copper anode was in one beaker with an NaCl solution, and the cathode was in the other beaker. As the current is applied, H2O is reduced to OH- and H2 at the anode and at the cathode Cl- is initially oxidized to Cl2. Ions move along the salt bridge to help maintain a balance of charges. The reason Cl2 is produced initially is because of a concept called overvoltage. If you do a search on google you'll be able to find a nice explanation about it. Basically, initially it is more favorable for the Cl- to be oxidized to Cl2 than it is for the OH- ions in solution to be oxidized to oxygen gas.

[jowrose]

I love how chemistry NEVER produces the results I think it will... I have much to learn, methinks.

 

Or maybe i'm stupid

[xeluc]

Yeah.. I definatly thought the exact same thing as you Jowrose. You'll find that after you learn about all the stuff you predicted wrong, it eventually gets a lot easier.

 

Now Jdurg.... I'm VERY confused here. In all my experience, hydrogen gas was given off at the cathode... Take this half reaction:

2H2O + 2e- ---> H2(g) + 2OH-

SO adding electrons (from the cathode) is making H2.

So either you are mistaken, or I am REALLY missing something here.

 

Also, "at the cathode Cl- is initially oxidized to Cl2". I have NEVER learned that oxidation was carried out at the cathode..

[jdurg]

My bad. I always get those two switched and forget the simple phrase 'An Ox, Red Cat'. Reduction is at the cathode and oxidation is at the anode. Let me go and edit that post.....

 

Heh. Too late. Can't edit.

[jowrose]

I tried the copper electrolysis with the salt bridge, and now I know what you mean by BLUE copper hydroxide... It's better than the mess of brown sludge I would always produce.

[xeluc]

I'm glad everything worked out for you Jowrose. Any time I have conducted electrolysis in NaCl solution without a salt bridge and a copper anode, I always got orange Cu2O. This might be what you are seeing also. SO anyway, now that I know I'm not going mad, how can you explain the hydroxide at the anode? If ANYTHING, the solution should become acidic if any Oxygen decided to become oxidized...

[ignilc]

I'm electrolyzing a solution of KNO3 with copper anode and stainless steel cathode.

Right when I turned the power on, the cathode produced lots of gas as normal, and the anode started corroding to a nice blue precipitate, but after about 2 minutes of runtime the SS cathode stopped bubbling completely, and the solution became filled with different hues of green and blue precipitates, also black. I also noticed that the cathode turned a slight copper colour at the top. My electrolyte is about 20 grams of KNO3 in 200ml of water. The cell draws a large current, I measured 2.8A at 5 volts.

 

The solution has a strong smell of ammonia.

 

What's going on? Why did the cathode stop producing hydrogen? Why did copper plate on the cathode? I thought Cu(OH)2 is insoluble. thanks

Edited May 31, 2010 by ignilc