Sciencemadness Discussion Board - Copper hydroxide synthesis question

Copper hydroxide synthesis question

Ramium - 3-12-2014 at 22:55

I am trying to make copper hydroxide.
In the electrolysis method i saw on youtube, they used a 16 v 3 amp power pack. What is the minimum voltage needed for this to work ( eg could i use a 12 v battery?)

[Edited on 4-12-2014 by Ramium]

Bert - 3-12-2014 at 23:20

Welcome! Take a look at the FAQ and please follow the posting guidelines there

Search on your topic first, then post questions in an existing related thread if possible. Or failing that, provide links to or copy of any references, stoichiometry of reactions. Document these and any other work you may have done so far-

gdflp - 4-12-2014 at 09:32

http://www.sciencemadness.org/talk/viewthread.php?tid=24174 This might be useful for you. To answer your question directly, yes a 12V battery would provide a high enough voltage.

Ramium - 4-12-2014 at 10:40

Thank you!

Ramium - 21-1-2015 at 20:09

I made copper hydroxide by adding a solution of NAOH to a solution of CUSO4 i dried the pricipitated Copper hydroxide in the sun but it turned to copper oxide. I read somewhere that copper hydroxide when heated turns to copper oxide so i thought when i try the experiment again I could evaporate it in a cold place and it shouldent turn to copper oxide is this so???????????????

woelen - 22-1-2015 at 00:00

A much lower voltage than 12 V is better. Use 5 V. That will give better results. With 12 V you could put 2 cells in series in order to have double speed. With 16 V you even could put 3 cells in series. Use bigger electrodes if you use cells in series in order to keep the resistance sufficiently low.

Copper hydroxide is only marginally stable. It very easily loses water and then it forms CuO. Even when stored under water, it can lose water. In practice, I would forget about copper hydroxide. Isolating the dry pure material is very difficult and it certainly will turn dark due to partial loss of water. Even when stored at room temperature, it will slowly turn to copper oxide.

So, you could also turn your project into making pure copper oxide. If you want to do that, then make copper hydroxide, suspended in water and rinse this to get rid of all kinds of other ions and then heat, while stiull under water. The precipitate then turns black and becomes more coarse and easier to separate from the water. You should even boil it some time. This makes the precipitate more compact and the boiling also assures that more other ions are leached out of the solid material. Then finally, you can filter and dry the black material.

This black material is a great source for other copper compounds. It dissolves easily in nearly every acid. Copper(II) oxide is one of the few metal oxides which does not become inert on heating and on storage.

[Edited on 22-1-15 by woelen]

Ramium - 22-1-2015 at 14:43

The electrolysis product seems quite stable but the predipitate product sounds unstable from what you described but it's the same chemical isn't it? So how can they be different in stability??????

Ramium - 22-1-2015 at 15:09

The copper hydroxide in this video is quite stable as it is not turning even slightly black

http://youtu.be/YsZodLOJnX8

Amos - 23-1-2015 at 10:22

I've found that the precipitate from using sodium hydroxide and a copper salt is very unstable, but if aqueous ammonia is used instead with an excess of copper sulfate, the product is stable pretty much indefinitely and is quite thermally stable too.

DraconicAcid - 23-1-2015 at 10:59

I have a big jar of it that the college bought from Sigma or Aldrich. It seems perfectly stable as a dry powder.

Ramium - 25-1-2015 at 21:29

Could you tell me how i could make copper hydroxide without it decomposing and with out electrolysis????

Ramium - 25-1-2015 at 21:36

Quote: Originally posted by No Tears Only Dreams Now

I've found that the precipitate from using sodium hydroxide and a copper salt is very unstable, but if aqueous ammonia is used instead with an excess of copper sulfate, the product is stable pretty much indefinitely and is quite thermally stable too.

Aqueous ammonia used instead of what?

DraconicAcid - 25-1-2015 at 21:38

Quote: Originally posted by Ramium

Quote: Originally posted by No Tears Only Dreams Now

Aqueous ammonia used instead of what?

Instead of sodium hydroxide. Just don't add an excess, or it will dissolve as a complex ion.

Amos - 26-1-2015 at 06:25

Quote: Originally posted by DraconicAcid

Quote: Originally posted by Ramium

Quote: Originally posted by No Tears Only Dreams Now

Aqueous ammonia used instead of what?

Instead of sodium hydroxide. Just don't add an excess, or it will dissolve as a complex ion.

I must also add that copper(II) chloride should not be used as the copper salt, or copper oxychloride will also form. Best to stick with good ol' copper sulfate.

blogfast25 - 26-1-2015 at 09:00

Quote: Originally posted by DraconicAcid

I have a big jar of it that the college bought from Sigma or Aldrich. It seems perfectly stable as a dry powder.

Freeze or vac dried, I'm guessing.

Amos - 26-1-2015 at 09:05

Quote: Originally posted by blogfast25

Quote: Originally posted by DraconicAcid

I have a big jar of it that the college bought from Sigma or Aldrich. It seems perfectly stable as a dry powder.

Freeze or vac dried, I'm guessing.

Not necessarily. Like I said, the kind I make with ammonia is light blue and very easily dries under a fan at 90C, or more slowly in air.

[Edited on 1-26-2015 by No Tears Only Dreams Now]

blogfast25 - 26-1-2015 at 10:09

Quote: Originally posted by No Tears Only Dreams Now

Not necessarily. Like I said, the kind I make with ammonia is light blue and very easily dries under a fan at 90C, or more slowly in air.

Can you describe your experimental procedure more accurately? Have you determined your product's composition?

[Edited on 26-1-2015 by blogfast25]

Ramium - 26-1-2015 at 12:33

Quote: Originally posted by No Tears Only Dreams Now

What are the quanties of ammonia and copper sulphate?

blogfast25 - 26-1-2015 at 13:12

Quote: Originally posted by Ramium

What are the quanties of ammonia and copper sulphate?

If you use fairly dilute NH3 and make sure not to exceed the NH3 stoichiometry of:

Cu2+ + 2 NH3 + 2 H2O === > Cu(OH)2(s) + 2 NH4+

... then only cupric hydroxide will form.

But with strong ammonia in excess then:

Cu(OH)2(s) + 4 NH3 === > [Cu(NH3)4]2+(aq) + 2 OH-

... will occur.

So keeping the molar ratio NH3/Cu2+ below 2 should result in Cu(OH)2 only.

Ramium - 26-1-2015 at 16:18

I brought some Ammonia at the supermarket i looked on the back and it said it was ammonium hydroxide could i use ammonium hydroxide instead of ammonia???????????

Sorry if this is a stupid question

[Edited on 27-1-2015 by Ramium]

blogfast25 - 26-1-2015 at 16:30

Quote: Originally posted by Ramium

I brought some Ammonia at the supermarket i looked on the back and it said it was ammonium hydroxide could i use ammonium hydroxide instead of ammonia???????????

Sorry if this is a stupid question

[Edited on 27-1-2015 by Ramium]

There is no such thing as 'ammonium hydroxide', that is an old and obsolete term you'll still find (but increasingly rarely) in some technical documentation.

What you bought was a solution of ammonia (NH3

in water (you did not specify any concentration although the bottle normally will).

When ammonia is dissolved in water the following chemical equilibrium is established:

NH3(aq) + H2O(l) < === > NH4+(aq) + OH-(aq)

... because ammonia is a weak alkali. Most ammonia in the solution is present as actual NH3, only a very small proportion of it is present as NH4+(aq) + OH-(aq).

[Edited on 27-1-2015 by blogfast25]

Ramium - 26-1-2015 at 18:50

Its 18 grams ammonia per litre of water. Would it work in the experiment?????

[Edited on 27-1-2015 by Ramium]

blogfast25 - 26-1-2015 at 19:19

Quote: Originally posted by Ramium

Its 18 grams ammonium hydroxide per litre. Would it work in the experiment?????

Tell me EXACTLY what it says on the label. Taking the value you cite at face value would mean a NH3 molarity of about 0.5 mol NH3 / litre of solution. That's very low for a commercial NH3 solution (but not impossible)

Describe the experiment you are planning to conduct.

[Edited on 27-1-2015 by blogfast25]

Ramium - 26-1-2015 at 19:57

It says ammonium hydroxide approximately 18/L.
I plan react it with copper sulphate to produce copper hydroxide
So would it work in the experiment??????

Amos - 26-1-2015 at 20:05

Yes, that would work, and be sure to use stirring while the ammonia is being added(you must add ammonia to copper sulfate, not the other way around) to keep any complex from forming. You should use small amount less ammonia(you'll need to do your own calculations there) than would be needed to fully react with the copper sulfate, leaving behind a dilute solution of copper sulfate on top of your precipitate.

Ramium - 26-1-2015 at 20:08

Thanks!

Ramium - 26-1-2015 at 23:00

Quote: Originally posted by blogfast25

Quote: Originally posted by Ramium

What are the quanties of ammonia and copper sulphate?

Thank you for this explanation. I now need to know how many grams of CuSO4 and how many ml of ammonia (which is 18gm/litre) i should use to make a decent amount of Cu(OH)2 (s). I would appreciate any help.

blogfast25 - 27-1-2015 at 05:43

For every 250 g of CuSO4.5H2O you need 70 g of 'ammonium hydroxide'.

Work it out from there and I will check your calculations.

Ramium - 27-1-2015 at 22:39

Quote: Originally posted by blogfast25

For every 250 g of CuSO4.5H2O you need 70 g of 'ammonium hydroxide'.

Work it out from there and I will check your calculations.

So if I had 250 gm of the CuSO4.5H2O
Would i weigh out 70gm of ' ammonium hydroxide' ?
Or if it says on bottle that it has 18g/l would that be 70g divided by 18 = 3.88 litres?

If it was 3.88 litres i would want to scale down those quantities

CHRIS25 - 28-1-2015 at 04:18

I will give you an example and maybe you could work out your own calculation? A standard bottle of Ammonia gas in water (some call it ammonium hydroxide but won't get into that), is 28%. You appear to have only 1.8% (18g/L). So when I do this to convert grams of gas per litre to mL liquid this is my calculation:
I need say 60 g of NH3 in my reaction.
I look at a chart that tells me Ammoniia solution has a Molecular weight of 35 g/mole (I am rounding everything off for convenience). This chart also tells me that this is 28% and has a specific gravity of 0.88 (leaving aside temperature fluctuations as this is not so important for most of us). So I need to do some maths:

60 g / 0.88 = 68 g Then 100% / 28% = 4
4 x 68 g = 243
This 243 is the number of mL that will contain around 60 g of ammonia gas.

[Edited on 28-1-2015 by CHRIS25]

blogfast25 - 28-1-2015 at 06:56

Quote: Originally posted by Ramium

Yes, 3.88 L for 250 g of CuSO4.5H2O.

Bear also in mind that 250 g of CuSO4.5H20 would have to be dissolved in about 1 L of water.

So scale down: 25 g of CuSO4.5H20 in 100 mL of water. Slowly and with constant intense stirring add 388 mL of your ammonia solution.

blogfast25 - 28-1-2015 at 07:02

Quote: Originally posted by CHRIS25

I look at a chart that tells me Ammoniia solution has a Molecular weight of 35 g/mole

Nope. An ammonia solution doesn't really have a molecular weight ('molar mass' or MM is a better term) because it's a solution comprised of 2 components: ammonia (NH3) and water (H2O). The MM of ammonia is 18 g/mol.

35 g/mol is the MM of the hypothetical 'ammonium hydroxide', NH4OH.

[Edited on 28-1-2015 by blogfast25]

CHRIS25 - 28-1-2015 at 07:11

Quote: Originally posted by blogfast25

Quote: Originally posted by CHRIS25

I look at a chart that tells me Ammoniia solution has a Molecular weight of 35 g/mole

I know that ammonia gas dissolved in water is not ammonium hydroxide, but I never ever took the weight of 17g/mol in any of my reactions because I was not obviously using the gas per se, using the solution I always used the 35g/mol. So are you saying that we should never use this latter figure?

[Edited on 28-1-2015 by CHRIS25]

[Edited on 28-1-2015 by CHRIS25]

blogfast25 - 28-1-2015 at 07:27

Quote: Originally posted by CHRIS25

So are you sating that we should never use this latter figure?

No, not really but you are at risk of confusing things.

Take the precipitation reaction we're talking about here. It can be written in two equivalent ways:

Cu2+ + 2 NH3 + 2 H2O === > Cu(OH)2(s) + 2 NH4+

or:

Cu2+ + 2 NH4OH === > Cu(OH)2(s) + 2 NH4+

In the first case you would need 2 x 18 g of NH3 per mol of CuSO4, in the second case 2 x 35 g of NH4OH per mol of CuSO4. Needless to say, these amounts are perfectly equivalent to each other.

If your ammonia solution, as is the general case, is expressed in % NH3 then it makes more sense to do the calculation in g NH3 solution needed to conduct your specific reaction.

[Edited on 28-1-2015 by blogfast25]

Ramium - 29-1-2015 at 00:15

Quote: Originally posted by blogfast25

Quote: Originally posted by Ramium

Thank you for that. I tried the experiment and ended up with a dark blue solution and a small amount of blue precipitate. I'm guess the precipitate is copper hydroxide and the solution is a copper complex. Is that so?

blogfast25 - 29-1-2015 at 08:17

Quote: Originally posted by Ramium

This indicates that you probably added too much ammonia solution. The dark blue is indeed the copper tetrammine - [Cu(NH3)4]2+ - complex.

The precipitate is Cu(OH)2.

The colour of the solution is deceptive though because the colour of the complex is very, very intense and so low concentrations of it can cause significant blueing of a solution.

Try again, reducing the amount of ammonia used by 20 to 40 %.

[Edited on 29-1-2015 by blogfast25]

gdflp - 29-1-2015 at 08:18

Yes, the blue precipitate is copper hydroxide. It sounds like you may have used too much ammonia, copper hydroxide is soluble in ammonia as it forms the water soluble complex [CuNH_3₄H₂O_₂](OH)₂, also known as Schweizer's Reagent. If you decant off the deep blue solution, and add a dilute acid to it slowly until you see a precipitate forming, you might be able to recover some more copper hydroxide.

blogfast25 - 29-1-2015 at 09:13

Quote: Originally posted by gdflp

[CuNH_3₄H₂O_₂](OH)₂, [...]

[CuNH_3₄H₂O_₂](OH)₂, also known as:

[Cu(NH3

4(H2O)2](OH)2

But adding HCl will not work here.

By adding HCl to [Cu(NH3

4(H2O)2](OH)2 the strongest base, the 2 OH-, would be neutralised first to [Cu(NH3

4(H2O)2]Cl2.

Adding further HCl will convert that to CuCl2 plus NH4Cl.

[Edited on 29-1-2015 by blogfast25]

Ramium - 1-2-2015 at 10:23

I tried again with less ammonia and got a much higher yield

blogfast25 - 1-2-2015 at 10:28

Well, there you go then!