[jaylan77 (OP)]

i performed electrolisis on a sodium chloride solusion for half an hour and was left with a rusty coulored sludge at the bottom
what is it?

[Chemistry4me]

Could be copper (I) oxide.

[damocles]

The electrolysis of brine can proceed in several different ways, depending on the nature of the electrodes, and the concentration of the sodium chloride.

With inert electrodes (silver, platinum, glassy carbon), you will produce hydrogen gas at the cathode, and either chlorine gas (concentrated NaCl) or oxygen gas (dilute NaCl) at the anode. There is a fairly fine balance between the evolution of these two gases.

Anode reactions: 2 H₂O --> 4 H⁺ + O₂ + 4 e^-
or    2 Cl^- --> Cl₂ + 2 e^-

But copper electrodes are far from inert. Instead of producing either of these gases at the anode, the most likely anode reaction is erosion (gradual dissolving) of the copper anode.

Anode reactions: Cu (solid) --> Cu²⁺ (solution) + 2 e^-  in dilute NaCl solution
or    Cu (solid) + 2 Cl^- --> CuCl₂^- + e^- in conc solution

As soon as there is a build up of copper ions in solution, the evolution of hydrogen gas at the cathode will stop, and be replaced with the reverse of either of the above reactions as a cathode reaction. The solid copper that is formed is unlikely to stick to the copper cathode unless it has a particularly clean surface, and no hydrogen clinging to it. It will far more likely be precipitated as a brownish sludge of copper metal.

Copper (I) oxide, Cu₂O, is a possible but most unlikely product of a cathode reaction. It is usually a bright cherry red colour rather than the typical rusty colour of metallic copper.

I am quite confident that your sludge is mostly, perhaps all, copper metal, and that your anode will disappear if you keep up the electrolysis for long.

[Bored chemist]

I have definitely seen Cu2O from electrolysis of salt water using Cu electrodes.
Electrolytically deposited copper is, in my observation, generally black (like most finely powdered metals).

[lightarrow]

i performed electrolisis on a sodium chloride solusion for half an hour and was left with a rusty coulored sludge at the bottom
what is it?

Maybe rust? 
You didn't write what the electrodes are made of. If it's iron or steel...

[damocles]

I have definitely seen Cu2O from electrolysis of salt water using Cu electrodes.
Electrolytically deposited copper is, in my observation, generally black (like most finely powdered metals).

(1) Interesting! Anode or cathode? Can you write a plausible electrode half-reaction (remembering that copper(I) chloride is insoluble and white, and that dichlorocuprate(I) is a yellow complex with a very strong association constant)?
 
[:I] Forget the first part -- it is obvious.

2 Cu + H₂O --> Cu₂O + 2 H⁺ + 2 e^- at the anode

That is quite a likely initial anode reaction -- would compete more than favourably with chlorine or oxygen evolution.

If the sludge is indeed Cu₂O , that would mean that at least some of it would need to flake off the anode before it could be further oxidized to Cu(II) species there. It would also necessarily mean a bright red anode during and after the electrolysis. Did the anode turn red?

(2) I agree that electrodeposited copper metal looks black on an electrode, but sludge that is formed by metal that fails to adhere is quite another matter.

------
For the original questioner: We have come up with two different plausible and likely explanations for your sludge. Can you make observations on the following:

a) Did the anode (positive electrode in your setup) noticeably erode?
b) Did the anode turn red?
c) Did the cathode (negative electrode) turn black?
d) From which electrode did the sludge seem to originate?
------

From Lightarrow:

Maybe rust? 
You didn't write what the electrodes are made of. If it's iron or steel...

It is true that this was not specified in the text of the original question, but check out the thread title ...
 []

[lightarrow]

With the font size I have, I could only read:
"What are the products of electrolysis of brine".

I don't have anylonger the sight I had once   :-)

About Cu₂O, there are several ways to test it. One of them consists in taking it (after washing several times with water) and adding an excess of concentrated HCl (in case heating gently): it goes in solution as dichloro copper complex.

[damocles]

About Cu₂O, there are several ways to test it. One of them consists in taking it (after washing several times with water) and adding an excess of concentrated HCl (in case heating gently): it goes in solution as dichloro copper complex.

Unfortunately, lightarrow, Cu metal will react exactly the same way with warm concentrated hydrochloric acid. A stream of hydrogen gas will slowly be evolved, but this will not be easy to notice.

I know that a lot of textbooks, especially the elementary ones, will say that copper does not react with concentrated hydrochloric acid, but try it and you will see.

[lightarrow]

Unfortunately, lightarrow, Cu metal will react exactly the same way with warm concentrated hydrochloric acid. A stream of hydrogen gas will slowly be evolved, but this will not be easy to notice.

I know that a lot of textbooks, especially the elementary ones, will say that copper does not react with concentrated hydrochloric acid, but try it and you will see.

Apart from  textbooks, our professor of General Chemistry at university (Chemistry faculty) used answers like yours to reject students immediately 
Copper (pure) cannot do that. If you are sure that there is copper and that hydrogen evolved, than it must be an alloy, for example brass or else.

[Bored chemist]

It would be rather difficult to check if conc HCl dissolved copper because you would need to remove dissolved oxygen first.
Anyway, another way of checking would be to add ammonia which will disolve half the Cu2O and turn the other half into Copper powder.
It might be interesting to add a saturated solution of salt and see if it's converted to CuCl.

However I don't agree with the initial supposed reason why it might be copper.
The anode mud in electrolytic refining cells isn't copper.

[damocles]

It would be rather difficult to check if conc HCl dissolved copper because you would need to remove dissolved oxygen first.
Anyway, another way of checking would be to add ammonia which will disolve half the Cu2O and turn the other half into Copper powder.
It might be interesting to add a saturated solution of salt and see if it's converted to CuCl.

However I don't agree with the initial supposed reason why it might be copper.
The anode mud in electrolytic refining cells isn't copper.

duhhh!

1. I was talking about a practical reaction between concentrated hydrochloric acid and copper in a normal laboratory setting, not an idealized reaction between ultrapure deoxygenated hydrocloric acid and copper in an inert atmosphere!

2. If you check back you will find that the suggestion I made was nothing to do with "anode mud". I was clearly referring to copper that was produced at the cathode and failing to adhere.

------

Lightarrow I have on many occasions taken copper foil or copper powder supplied as laboratory reagents and added laboratory concentrated hydrochloric acid and seen a yellow coloration slowly develop in the acid, and occasionally streaming of minute gas bubbles (the latter observation, incidentally, suggests that dissolved oxygen is not involved).
I invite you to do likewise. I doubt that many of my colleagues have ever done so, and if they have not done the experiment themselves, they tend to take textbook statements as authoritative. The yellow coloration, incidentally, may not be due to CuCl₂^–, which "should" be colorless, but to the product of its further oxidation CuCl₄^2–, which certainly is yellow.

Why have I done this experiment so many times? In my honours year project I had occasion to add concentrated hydrochloric acid to 3-hydroxyisothiazole that was in a copper foil container. When I observed the yellow coloration, I had to check what was going on, and as part of that I just put the acid with the foil and no organic substrate, and the yellow coloration was again observed. End of story as far as my project was concerned.

But I was intrigued by this apparent reaction that contradicted the textbooks, and have tried it again since with several grades of copper in laboratories in both Australia and England. I have never tried to publish. I am not asking you to accept my authority -- I am merely inviting you to try a very simple experiment for yourself.

[damocles]

Further to the reaction between copper and concentrated hydrochloric acid:

I found the following web reference to a peer reviewed article by Swedish nuclear scientists (chemical engineers or chemists):
http://www.stralsakerhetsmyndigheten.se/Global/Publikationer/SKI_import/010803/04318226039/98-19.pdf

In it I could find the following delta G°_f values:
CuCl₂^– = –246.0 kJ/mol
Cl^– = -131.2 kJ/mol

therefore in the reaction Cu + H⁺ + 2Cl^– --> CuCl₂^– + ¹/₂ H₂

delta G°_r is +16.4 kJ/mol

K_eq is therefore: exp{–16.4/(8.314 * 0.298)} = 0.00133

Concentrated HCl as supplied is above 10M in H⁺ and Cl^–

K_eq = [CuCl₂^–]/([H⁺].[Cl^–]²)

In 10 M HCl solution [CuCl₂^–]_eq = 0.00133 x 1000 = 1.33 M

Clearly this reaction can proceed to a significant extent!

(My apologies to forum readers who are not trained chemists, but there was a point here that needed to be addressed).

[lightarrow]

Further to the reaction between copper and concentrated hydrochloric acid:

I found the following web reference to a peer reviewed article by Swedish nuclear scientists (chemical engineers or chemists):
http://www.stralsakerhetsmyndigheten.se/Global/Publikationer/SKI_import/010803/04318226039/98-19.pdf

In it I could find the following delta G°_f values:
CuCl₂^– = –246.0 kJ/mol
Cl^– = -131.2 kJ/mol

therefore in the reaction Cu + H⁺ + 2Cl^– --> CuCl₂^– + ¹/₂ H₂

delta G°_r is +16.4 kJ/mol

K_eq is therefore: exp{–16.4/(8.314 * 0.298)} = 0.00133

Concentrated HCl as supplied is above 10M in H⁺ and Cl^–

K_eq = [CuCl₂^–]/([H⁺].[Cl^–]²)

In 10 M HCl solution [CuCl₂^–]_eq = 0.00133 x 1000 = 1.33 M

Clearly this reaction can proceed to a significant extent!

(My apologies to forum readers who are not trained chemists, but there was a point here that needed to be addressed).

10M HCl is not totally dissociated but this should be irrelevant for the order of magnitude of the result. However the activity coefficients should be quite significantly different from one at such high concentrations.
Anyway, I will try the experiment with my poor means when I'll have time.
Bye!

--
lightarrow