Copper II Chloride, Sulfate; and Nails

Name: Roy
Status: educator
Grade: 9-12
Location:

Question: I was performing a laboratory exercise with my students.
My original plan which I have done before was to place two nails in
a copper II chloride solution (1 molar) watch the copper precipitate
on the nails, wash it off and dry it before weighing it, calculating
yield etc. However, I could not find any copper II chloride, so I
substituted copper II sulfate. I did not have enough to make a 1 M
solution so I made 0.7 thinking it would be OK. What we observed was
somewhat different than I had observed with copper II chloride.
Rather than copper "fuzz" forming on the nails, it was more like
copper plate. The plating would flake off, but it seemed stuck to
the nail tighter than the "fuzz". I do not think the yield will be
as high as I have seen in the past. Another thing that was different
was that the solution did not change color from blue to green, it
stayed blue. Why would it not work with the sulfate as well as it
did with the chloride? They are just spectator ions, are they not? I
may run this again in a couple of weeks. Should I order the copper
chloride or make a change to the procedure to get a better yield
with copper sulfate?
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Roy-
  I would order the chloride, if only because you are out of CuSO4.
But meanwhile try adding salt to your copper sulfate solution.
It might help enough.

  Unfortunately they are not entirely spectator ions.
They can speed up or slow down the reaction.
Chloride helps by complexing with the dissolved Iron ions,
and even more by catalyzing dissolution of iron oxides
(chloride is famous for catalyzing many metal-oxide transformations;
including salt-water -> rust).
Sulfate might well hinder fast corrosion a little
simply because iron sulfate is only slightly soluble,
so a thick wet crust will form that slows access to the iron surface.

In electroplating, going slow often makes dense, shiny, hard metal films,
while going too fast degenerates into dark rough plating or even "fuzz".
Slowing by sulfate crust might be responsible for the denser copper plate you were getting.
Your less-concentrated 0.7M solution is also likely to have a slowing effect,
but I not as much slowing as you've observed.

If your students have access to a scale sensitive enough to measure
the lesser masses of copper plating and iron dissolution that happen,
your sulfate alone might be adequate for a class experiment.

Adding a roughly equal concentration of NaCl might speed it up enough
for your class needs.

If instead of NaCl you add CaCl2,
maybe much of the sulfate will precipitate out as CaSO4.
Then you could filter it out.
Then even the iron-sulfate crusts wouldn't get in the way.

Solubilities:
CaSO4 ~ 3gm/L,  CuSO4 ~ 140gm/L,
CaCl2 ~750gm/L, CuCl2 ~700gm/L,
FeCL2, ~700, FeCl3 ~700,
Fe(II) and Fe(III)  sulfates, variable.  More with Cl- around.

Another thing: Cu chloride can be in two oxidation states:
Cu(II): Cu+2: CuCl2,   or
Cu(I): Cu+1: CuCl.
Because Cl- helps stabilize Cu+1 ions by complexing with them in solution.
Going from Cu+2 to Cu+1 is a free oxidation source for the iron;
no complementary precipitation of Cu need occur at first corrosion of Fe.
Corrosion gets a big head start and the surface gets rough and hydrated,
difficult for precipitates to encapsulate later.

Sulfate cannot complex the Cu+1 ion like chloride does.
It seems only Cu+2 and Cu metal can exist in sulfate solution.
So right from the beginning the dissolving iron has to
struggle its way out through a smothering blanket of precipitating copper.
It is slower.

Jim Swenson
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   There are a number of differences that could affect the outcome of "sulfate" 
   vs. "chloride".
1. The anion is not just a spectator. The morphology of the electrolysis can be 
very sensitive to the choice of counter ion.
2. You changed the concentration of the solution. This by itself could change 
the growth of copper metal.
Among other things the electrical conductivity between the two salts could have 
an effect. Stirring vs. diffusion could also have played a role.
3. There might be different impurities, differences in pH (that you did not 
mention), in addition to 1. and 2. above.
   Do not give up!! You may have the makings of a very interesting science project. 
   A few possible variations that could yield some differences worth noting.
   A. Adjust the pH with HCl or H2SO4 and compare the "yield" as well as the quality 
   of the metal deposition.
   B. Observe any differences in the nails. Not all nails are created equal. There 
   could be differences in the composition of the nails you are using that dictate 
   the different observations you made.
   C. Change the surface tension of the solution using "dish" soap compared to 
   "pure" water.
   D. Change the viscosity of the solution with a "thickener" -- starch, 
   polyethylene oxide, or other water soluble polymer to observe changes compared 
   to the "standard".
   E. Do the experiment with an excess of ammonia that forms the ammonia complex 
   with copper.
This list could get as long as your patience. I do not know the answer to the 
differences you might observe, but that is the nature of research. If you know 
the answer in advance, why do the experiment?

Vince Calder
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