Sciencemadness Discussion Board - Addition of sodium bromide to copper sulfate solution

Addition of sodium bromide to copper sulfate solution

sternman318 - 2-6-2011 at 12:11

I have some unexplained chemistry...

When you add sodium bromide crystals to a solution( ive only tested it with somewhat dilute solutions), a dark brownish-violet solution forms and lingers around them, but dissolves and loses its color when the solution is stirred. Even after adding copious amounts of NaBr, the brown color appears, then completely dissappears when stirred. The solution appears somewhat greenish, almost like that of copper chloride. I wish to do a few more experiements

Add a NaBr solution to a CuSO4 solution
Repeat, but with an acidified CuSO4 solution
Add CuSO4 crystals to a NaBr solution

Butttttt, I nearly wasted all of my NaBr in the process of the initial experiment ( time to make a run to Walmart haha).

Does anyone know what is occuring? Could I be forming elemental bromine, in a reaction analagous to that of iodine and a copper solution?

Jor - 2-6-2011 at 12:17

In very concentrated bromide solutions, copper forms a dark purple/violet bromocomplex, but this is only at very high concentration. For example, when you add some copper sulfate crystals to 40% HBr, you get a dark violet solution, wich however turns green when you dilute it.

sternman318 - 5-6-2011 at 10:08

New invisible ink?
http://www.youtube.com/watch?v=7WlGIQUzvUA

Excuse the use of a toilet and my inability to speak ( no sleep). And note the solution is as it appears in the video, I am not sure why I called it blue.

Does anyone know what the formula is, specifically? I think I am going to try to get some crystals out of the remaining solution

LanthanumK - 7-6-2011 at 04:38

It does appear to be a bromine complex with the copper.

My Wal-mart does not have sodium bromide in it.

[Edited on 7-6-2011 by LanthanumK]

m1tanker78 - 7-6-2011 at 08:00

Quote: Originally posted by LanthanumK

It does appear to be a bromine complex with the copper.

My Wal-mart does not have sodium bromide in it.

[Edited on 7-6-2011 by LanthanumK]

Lowe's has it in the pool chemical section (outside/garden). It's sold in small packets; the larger containers have some organic bromo compound (can't remember ATM).

Tank

sternman318 - 7-6-2011 at 17:31

Yup, pool section is where it's at! Apparently people use bromine as an alternative to chlorine in their pools. If only they sold sodium iodide -_-

entropy51 - 7-6-2011 at 17:59

Quote: Originally posted by sternman318

If only they sold sodium iodide -_-

Well, they do, sort of.

LanthanumK - 8-6-2011 at 03:59

Reduction of NaI*I2 (tincture of iodine) may produce NaI. Around here, 1 fl oz of NaI*I2 is readily available.

sternman318 - 8-6-2011 at 07:49

Quote: Originally posted by LanthanumK

Reduction of NaI*I2 (tincture of iodine) may produce NaI. Around here, 1 fl oz of NaI*I2 is readily available.

Really? I've been looking for that- where do you get it? I have only seen to have been able to find povidone-iodine solution, but have not bothered trying to work with that because I really wouldn't know what I waqs doing haha. Though, I think I should just order some online, probably end up being cheaper.

LanthanumK - 8-6-2011 at 07:52

I bought it at Walgreens or CVS, in my area. It is about the same price online as in those stores.

sternman318 - 10-6-2011 at 12:22

So I have had my solution sitting out to slowly evaporate. Everyday it is getting darker, going from the nice pale green to a brownish color. however, when you cool in down it, surprisingly, goes back to the pale green color. Allowing it to return to room temperature yields the brown solution again.

This all seems quite confusing to me. However, looking at the wikipedia page for copper chloride...

"Aqueous solution prepared from copper(II) chloride contain a range of copper(II) complexes depending on concentration, temperature,and the presence of additional chloride ions."

I am going to go ahead and assume that this is the case, just with bromide instead. So my question is, how would I go about crystallizing this solution? even small amounts of it do not seem to dry up, and remain as dark brown drops. Should I be worried about heating the solution? I am somewhat apprehensive towards heating solutions that I am unfamiliar with, after trying to dehydrate a solution of NaNO3 and accidently evolving NO2, i believe.

[Edited on 10-6-2011 by sternman318]

Arthur Dent - 10-6-2011 at 13:03

Quote: Originally posted by sternman318

...when you cool in down it, surprisingly, goes back to the pale green color. Allowing it to return to room temperature yields the brown solution again...

Welcome to the magic of Copper chemistry. Try something cool, just add a bit of hydrogen peroxide to your brown solution, and it should turn a bright emerald green!

Robert

sternman318 - 12-6-2011 at 07:01

Robert- What should the above do? I dont want to waste my solution because I am trying to dry it out xD Or maybe I will just prepare some more later

But...

Copper is crazy, let me just say that.
I have accidently stumbled upon a method to create what I think is verdigris , but I dont think that that is very useful haha. But it is a very pretty seafoam.

I have a gorgeous blue/purple solution, but my poor lab practices means that I didnt record what I did to it, and have no idea what is in it

And after trying to make some crude tetraaminecopper ions by adding what should be ammonium nitrate and then neutralizing with sodium bicarbonate, I added some NaBr and it yielded a bright yellow precipitate. The precipitate dissolved when disturbed so I think that two things might be happening

1) that is the color of the solution/salt ( [Cu(NH3)4] Br2 ? or even [Cu(NH3)nBr4-n]

2) The copper is forming the [Cu(Br)4] 2-, and some ammonia is getting in there, making it yellow- which is pretty much the same thing as 1) haha.

If anyone would like to try and repeat the above using better chemicals and technique, I think the yellow color would persist if more bromide was added. I, for some reason, dont have test tubes, so I end up wasting a lot of reagents trying to do things.

[Edited on 12-6-2011 by sternman318]

m1tanker78 - 12-6-2011 at 07:43

Quote:

I, for some reason, dont have test tubes, so I end up wasting a lot of reagents trying to do things.

Buy some cheap shot glasses at walmart. I think a 6 pack costs 2 or 3 bucks. They're nice for small scale experiments and allow complete visibility - even fine precipitate that settles to the bottom and such. You can't heat stuff in them (such as over a burner), though!

Tank

sternman318 - 16-6-2011 at 16:42

Well I cooked my solution, and made some asphalt!

Or CuBr2

But if any of you are working on your copper rainbow, here is your dark dark purple/black copper!

And thanks for the suggestion Tank!

sternman318 - 16-6-2011 at 18:37

This was the color change when a pretty saturated solution was cooled ( furthest left) then allowed to warm to room temperature ( right)

[Edited on 17-6-2011 by sternman318]

woelen - 16-6-2011 at 22:44

Yes, copper is also one of my favorite elements. I have done a lot of experimenting with it. Combinations of copper with halide ions are particularly interesting. It also is worth experimenting in non-aqueous solvents. A funny thing is to dissolve some CuCl2 or some of your newly made CuBr2 (which indeed is black, I have a commercial sample which looks very much like yours) in acetone or methanol.

Reduction of copper(II) in strong HCl or strong HBr also gives very interesting results. Exceedingly intensely colored complexes are formed, which contain both copper(I) and copper(II). As reductor you can use plain copper wire, the copper reduces copper(II) to copper(I) and the metal also becomes copper(I).

Copper Bromide > intresting properties?

numos - 26-2-2014 at 16:50

So I've created quite a bit of Copper bromide from sodium bromide and copper II sulfate (both were concentrated solutions in water).

I wish to know how it functions. So at high concentrations the solution turns a beautiful maroon color, and when diluted turns green until it dilutes to clearness. What's interesting is the sudden color change, it seems almost immediate without a gradient, (Unless you cool it in which case its gradual change) although it's more likely that the gradient occurs at a very specific range over a few percentages, so more accurate tests will have to verify that.

I've set it out to evaporate and so far I got this [see image 1] I've taken the somewhat dry crystals (image 2) and put them in a desiccator for now as they don't seem to be air drying. Once they are completely dry I will run tests to see at what concentrations the color change occurs .

Any ideas of what actually happens? It's definitely not a chemical reaction that causes the color change right? Is this substance (the green and the brown) indeed copper bromide CuBr? This picture I found is copper bromide so why isn't the (almost) dried version of CuBr similar to this greenish color?

NOTE: The side product of this reaction is sodium sulfate and I did not recrystallize this product out as I am in high school, funding myself is not easy and I have yet to obtain a hotplate. Would this compound contribute to this anomaly?

Also notice in image 3, my poor spatula... It was perfectly stainless - and now it's corroded, as I was scooping out the Copper bromide form the evaporating pan this happened! It ate steel!

[Edited on 2-27-2014 by numos]

woelen - 27-2-2014 at 06:41

The red/brown color is due to the complex ion CuBr4(2-), which contains copper in oxidation state +2.

The first picture you show is impure CuBr, which contains copper in oxidation state +1. Pure CuBr is white, the green color is due to partial oxidation and this green material is CuBr with some copper(II) and oxide in it as well.

nezza - 27-2-2014 at 06:45

I suspect the colours are due to different copper complexes.

In dilute solution the hexaquo ion Cu(H2O)6 2+ ion dominates. This is the typical blue colour.

With chloride in concentrated solution the chloride complexes with copper giving Cu(Cl)4 2- ions. This gives a brown solution.

I suspect bromide acts similarly to chloride. Depending on the concentrations of the different ions many colours are possible.

Texium - 6-9-2014 at 19:42

Just wanted to verify an idea that I had. Would making a bromine generator that passes bromine vapor directly onto CuO be an effective way of making copper(II) bromide? It seems like that might be a cleaner way of producing it if it works. I thought that I would try it tomorrow, but I wanted to make sure that I wouldn't be completely wasting my time and making a huge mess by doing so.

Also, is manganese dioxide a good enough oxidizer to liberate bromine when used with sodium bromide and conc. sulfuric acid? I previously used potassium permanganate, but I am currently out of it.

[Edited on 9-7-2014 by zts16]

Zyklon-A - 6-9-2014 at 20:16

How would you ballance that equation?
CuO + Br₂ = CuBr₂ + 1/2 O₂
Seems unlikely to go forward very well. Would elemental copper work, or would the reaction be to slow?

Texium - 6-9-2014 at 20:30

Quote: Originally posted by Zyklon-A

How would you ballance that equation?
CuO + Br₂ = CuBr₂ + 1/2 O₂
Seems unlikely to go forward very well. Would elemental copper work, or would the reaction be to slow?

2CuO + 2Br₂ --> 2CuBr₂ + O₂

I read that elemental copper would work, but I thought it seemed like CuO might work better and react faster. Wikipedia says that it will react with hydrobromic acid, so I thought that it might react with elemental bromine too?
Another thing I was just thinking about, would it be better to allow the bromine to react as a gas, or should I condense it and let it drip onto the CuO?

Brain&Force - 6-9-2014 at 21:29

The problem is that HBr is an acid and CuO is a base so a neutralization takes place. Not so with bromine gas.

Texium - 6-9-2014 at 21:56

Quote: Originally posted by Brain&Force

The problem is that HBr is an acid and CuO is a base so a neutralization takes place. Not so with bromine gas.

Oh, of course! Should have realized that! (Late night chemistry planning = many careless mistakes)
Well, in that case, I'll make a hydrogen bromide generator instead of bromine. Using dilute sulfuric acid and no oxidizing agent should do the trick, and then the HBr gas produced can react with the CuO. I will try this tomorrow.

chornedsnorkack - 6-9-2014 at 23:27

Quote: Originally posted by Zyklon-A

How would you ballance that equation?
CuO + Br₂ = CuBr₂ + 1/2 O₂
Seems unlikely to go forward very well. Would elemental copper work, or would the reaction be to slow?

Copper obviously
Cu + Br₂ = CuBr₂

CuO? A few ways to balance it:
CuO + H₂O + 2 Br₂ = CuBr₂ + 2HBrO (CuO is a weak base, HBrO a weak acid, so copper hypobromite not expected)

6CuO + 6Br₂ = 5CuBr₂ + Cu(BrO₃)₂

chloric1 - 21-12-2022 at 20:18

Elemental bromine is nasty as hell and you need special respirator for it! One nice property of bromine is in when added to an excess of strong ammonia solution, it is taken up to form clear ammonium bromide solution and nitrogen gas via the reaction: 3 Br2 +2 NH3 > N2+ 6HBr. As you can see the need for excess ammonia as if not dangerous nitrogen tribromide could form! Anyways not only is this a safe way to neutralize bromine, it creates useful ammonium bromide which is de facto hydrobromic acid in solid form. You should be able to heat copper oxide with ammonium bromide to get cupric bromide.