This video will be another installment in the Elemental Syntheses series, this time
on copper.
I will demonstrate two methods of obtaining pure copper, and discuss to a small degree
how it is produced industrially.
We will then look at a few complexes of copper in various oxidation states, and the first
instance on YouTube of copper(III), that is, copper in the +3 oxidation state.
So, first off is using electrolysis to purify copper.
Two electrodes are prepared, one is the impure copper, for example a piece of copper pipe,
copper coins, etc. and the other a piece of pure copper.
The pure copper should be connected to the negative terminal which is called the cathode,
and the impure copper to the positive terminal called the anode.
The two electrodes are then placed into a solution of a copper salt, I used copper sulphate.
A current is then passed through the cell, and copper will slowly plate out from the
anode onto the cathode.
The ideal voltage to use is 0.15 to 0.3 volts, although higher voltages do also work.
The higher the current, the faster the electroplating will take place.
It works because the ions in solution are the only ions that can move from one electrode
to the other.
Since only copper ions are present, only copper will be transferred, and any other metals
will gradually drop off the electrode as a sludgy powder (in theory).
Industrially, this is done to purify the copper obtained from the furnace, known as "blister
copper".
The impurities are often rich in precious metals like silver, gold and platinum, as
well as relatively worthless ones, and slag from the furnace.
The purest copper one is likely to be able to find around the home is copper from electrical
wiring; pure copper is a much better conductor when pure, and is easier to form into wires
pure, so the extra step is taken to do this for the increased performance.
The copper in piping is much less pure, and has been alloyed to improve strength, corrosion
resistance and other desireable properties.
Just a little note: for those watching from the UK, any copper coin minted before 1992
is 97.5% copper, 2.0% zinc, and 0.5% tin.
For those watching from the US, any penny minted prior to 2008 is 95% copper, and 5%
zinc and tin.
Post these dates, the coin is just copper-plated steel and zinc respectively.
An alternative electrical method is to simply electrolyse a solution of a copper salt, and
copper will precipitate at the cathode (negative terminal), and leave an acid in solution.
Sulphuric acid from copper sulphate, hydrochloric from the chloride, etc.
A chemical approach to making pure copper is simple displacement from solution by using
a more reactive metal, like zinc, aluminium or magnesium.
A solution of the copper salt is prepared, and the more reactive metal added.
I used magnesium powder because I don't have any zinc powder, and I haven't had much success
with aluminium powder in the past.
I used powder simply to make the reaction take place more quickly.
Aluminium foil, for example, gives good results when doing this, but it takes a few days for
the reaction to complete.
Anyway, after all the magnesium has reacted, I added a little sulphuric acid to dissolve
any unreacted metal away, and then filtered the solution to reveal the reddish copper
powder.
It was washed with water, and allowed to dry on the filter-paper.
So, that's how copper metal can be prepared.
I shall now investigate a few of the complexes of copper in the +1 and +2 oxidation states,
before preparing some potassium cuprate(III).
The first ones I investigate are the ammine and chloride complexes of copper(I).
Cuprous ions are not very easy to prepare due to their tendency to either oxidise in
the air, or disproportionate into copper and cupric ions.
So I first make a solution of cupric and cuprous chlorides in hydrochloric acid (because I
have a mixture of the two available), and then add some sodium metabisulphite to reduce
the copper(II) ions to copper(I), which then complex with the chloride ions present in
solution to form the tetrachloridocuprate(I) complex, also written: [CuCl4]2-.
This ion is a yellow-brown colour, and this can be just about seen through the murky precipitate
that has formed.
I'm not sure what this precipitate is, but I'm sure the lack of purity has something
to do with it.
To prepare the ammine complex, I prepare the same solution, but then add concentrated ammonia
solution, and a different coloured complex forms, this time with a blue tint; though
I fear this may be copper(II) interfering.
Now for some copper(II) complexes.
I will prepare four of these, all starting from cupric sulphate.
A solution of copper sulphate is added to each of four test tubes, and then concentrated
ammonia solution is added to the first one, nothing to the second, sodium chloride to
the third, and potassium bromide to the last one.
The ammonia initially reacts as an alkali, and precipitates copper hydroxide, but this
then redissolves in the ammonia to form the tetraammine complex.
The complexes are, from left to right: [Cu(NH3)4(H2O)2]2+, [Cu(H2O)6]2+, [CuCl4(H2O)2]2-, [CuBr4(H2O)2]2-.
I will refer to these as the ammine complex, aqua complex, chloride complex, and bromide
complex.
I have placed them in a very careful order, and as you may be able to see they very roughly
fall along the spectrum(!).
Complexes form their characteristic colours due to them absorbing all but certain frequencies
of light, and reflecting the remainder.
So, the ammine complex absorbs light from red to green, but reflects blue light; so
we see a blue colour.
Different ligands cause different frequencies of light to be absorbed, and this is consistant
from metal to metal.
A spectroscopic series can then be constructed, and this is the order I have placed the test
tube in.
Unfortunately, I don't have very many usable ligands, hence the shortness of the row, but
this should demonstrate the principle.
Finally, I will show how I forced some copper into the +3 oxidation state, a very rare and
unusual one only seen commonly in superconductors.
I first prepared some oxygen as in my previous video, Elemental Syntheses - Oxygen, and then
used a lid which I had previously burned to remove any organic matter to seal the jar
of oxygen.
I then had to very carefully and quickly open the jar upside down (since oxygen is slightly
less dense than air), place a piece of potassium on the centre of the lid, cover it in copper
oxide powder, and reseal the jar.
The rusty lid I used had actually got a few tiny holes, so the next step wouldn't result
in a pressure explosion.
I heated the potassium covered in copper(II) oxide with a spirit burner until it caught
fire, continued heating and kept doing so for about 5 minutes after the potassium had
burned out to ensure reaction had taken place.
The idea of all this is to react copper oxide and potassium peroxide or superoxide under
and atmosphere of oxygen.
Since simple cuprates are so unstable, the oxygen atmosphere is required to ensure the
reaction product doesn't decompose too easily, and helps it form in the first place.
So after all that, did it work?
It did indeed, and I am very proud to present the first instance on YouTube of a simple
cuprate salt!
The blue section you see is the potassium cuprate(III), the black unreacted copper oxide,
and the yellowy bit is likely either potassium peroxide or superoxide.
Just as further proof that the blue was indeed copper(III), I left the lid out in open air
for a few hours, and I came back to a black sludge of copper(II) oxide.
It had decomposed from the moisture in the air to form cupric oxide and potassium hydroxide,
which absorbed more moisture, and so on until it had all decomposed.
So, that's all I've got for copper, thanks for watching, please subscribe, rate and comment!