Question about tetrabromocuprate ion

[latentheat]

I added an excess of sodium bromide to a solution of copper sulfate to produce this complex ion.

 

 

I think that's how it's supposed to look... it's a deep purple.

 

If I was to evaporate the water would I get a mixture of sodium sulfate and sodium tetrabromocuprate(II)?

 

Would it be possible to get one to crystallize before the other?

[latentheat]

Ok, decided to play with it a little. I heated it and just as the liquid begain to boil it got all weird lookin. I though I had just screwed everything up so I began to dilute it to get rid of it.

 

Then, as I was going to pour it away, I noticed a greenish/blueish powder had precipitated.

 

Copper bromide maybe? Any ideas?

[woelen]

Ok' date=' decided to play with it a little. I heated it and just as the liquid begain to boil it got all weird lookin. I though I had just screwed everything up so I began to dilute it to get rid of it.

 

Then, as I was going to pour it away, I noticed a greenish/blueish powder had precipitated.

 

Copper bromide maybe? Any ideas?[/quote']

If you did not have any excess acid, then the greenish/bluish powder most likely is basic copper sulfate, contaminated with basic copper bromide.

 

The composition of this compound is xCu(OH)2.yCuBr2.zCuSO4 of indeterminate stoichiometry.

 

Aqueous copper ions hydrolyse quite strongly (one of the ligand-water molecules splits off a H(+) ion):

 

[Cu(H2O)6](2+) <---> [Cu(H2O)5(OH)](+) + H(+)

 

With bromide (in fact with many anions, such as sulfate, chloride, acetate, etc.), a basic copper salt precipitates from solution.

 

You'll see a similar effect if you just take a solution of CuSO4.5H2O. If you let this stand for a long time, then you'll get some precipitate of basic copper sulfate and the liquid becomes somewhat acidic. On heating, this effect can be observed more quickly.

 

Many metals suffer from this effect. It is even stronger for iron (e.g. a solution of ferric sulfate or ferric chloride always is somewhat turbid, due to hydrolysis and precipitation of basic iron salts) and it is excessive for bismuth. Even in 1 M HNO3, solutions of bismuth salts precipitate basic salts, such as

Bi(OH)2 NO3.

[woelen]

I added an excess of sodium bromide to a solution of copper sulfate to produce this complex ion.

<images snipped' date=' see original post>

I think that's how it's supposed to look... it's a deep purple.

 

If I was to evaporate the water would I get a mixture of sodium sulfate and sodium tetrabromocuprate(II)?

 

Would it be possible to get one to crystallize before the other?[/quote']

You will get highly contaminated Na2SO4. Na2CuBr4 is very very hygroscopic. You'll end up with a kind of thick syrup, which on standing separates Na2SO4 and NaBr. Getting the Cu-salt in solid form hardly is possible. When you still continue heating, then besides water, you'll drive off impure HBr, leaving a basic mix behind. I think you'll get some brown/green mix, which is mostly insoluble in water, because it will be quite basic.

 

It is remarkably diffiicult to obtain such complex at high purity, due to the hygroscopic nature of them and due to the effect that they suffer severely from hydrolysis.

[budullewraagh]

i dont agree. Na2SO4 and NaBr? where do all the cupric cations go? they certainly are not reduced.

 

i definitely can see the Na2SO4 being formed, along with CuBr2. do you agree?

[latentheat]

Thanks for the detailed responses, woelen. I knew you would respond because you seem to do a lot of work with complexes. I love the experiments on your site with pictures. That way if we don't have the reagents to do the experiment ourselves we still get to see what happens.

 

So, just at of curiosity, because of the hydroscopic nature of Na2CuBr4, how is it prepared in industry? Is it prepared at all?

[woelen]

i dont agree. Na2SO4 and NaBr? where do all the cupric cations go? they certainly are not reduced.

 

i definitely can see the Na2SO4 being formed' date=' along with CuBr2. do you agree?[/quote']

If you evaporate away the water, then the concentration of all ions increases (same number of ions in smaller volume of liquid). In the brown liquid, shown by Latentheat, you have Na(+), SO4(2-), Br(-), CuBr4(2-). I also did this experiment and from that I remember that you need a very concentrated solution of NaBr in order to get the brown/purple complex, so there is excess Na(+) and Br(-). As the compound Na2CuBr4 is very soluble, the first to separate from the liquid are Na(+), Br(-) and SO4(2-). So, you will get a crud consisting of a mix of NaBr and Na2SO4, highly contaminated with copper species. Most of the copper remains in solution as CuBr4(2-).

 

Indeed, no copper is reduced.

[woelen]

Thanks for the detailed responses' date=' woelen. I knew you would respond because you seem to do a lot of work with complexes. I love the experiments on your site with pictures. That way if we don't have the reagents to do the experiment ourselves we still get to see what happens.

 

So, just at of curiosity, because of the hydroscopic nature of Na2CuBr4, how is it prepared in industry? Is it prepared at all?[/quote']

I doubt whether a compound like Na2CuBr4 is prepared in industry. What would be its commercial application?

Production of very hygroscopic compounds, which decompose at elevated temperatures, frequently involves drying under reduced pressure over H2SO4 or any other non-volatile strongly water-attracting agent. At the very low pressure, the water is removed more easily and the water is absorbed by the H2SO4

[akcapr]

complexes- wat seperates them from compounds; just the ligands and the central atom, or is there more?