Element Collection

Element Collection

Friday, May 20, 2011

Copper Compound Conundrum, Part 3

Copper Compound Conundrum Part 3: Testing the Precipitate - the exciting conclusion!

Now that I had determined the identity of the solution, it was time to figure out what the blue precipitate was. Remember from Part 2, this is what my mystery concoction looked like (in the beaker on the right).

I scooped some of the blue solid out of the bottom, placed it in a filter paper, and rinsed it with water to remove any of the green solution. Surprisingly, all the solids dissolved!

I ended up with this clear blue solution. My blue crystals going easily into solution like this lead me to believe that this must be the copper sulfate I was looking for. That was the theory, so time to test it! I transferred some solution into two test tubes.

I am holding a small piece of iron wire in front of the left tube - that will be test #1. If it is indeed copper sulfate, it should react with iron to displace copper metal:
CuSO4 + Fe == FeSO4 + Cu

To the right test tube, I will perform an anion test, similar to what I did in Part 2. This time I am looking for the sulfate ion, so I made up a solution of barium chloride. If sulfate is present, a white precipitate of barium sulfate will form. This is a photo of the two tests taking place:

We can see the piece of iron on the left is visibly reddish - this is due to copper plating out of solution as iron ions enter it. On the right, we see a white precipitate that immediately forms when the barium chloride solution is added, which looks just like Part 2's test.

Testing complete:

After a few minutes, quite a bit of copper has come out of solution and plated onto the wire in the left test tube. Notice also that the blue color of copper(II) ions has disappeared as iron(II) ions have replaced them. A white cloudiness of suspended barium sulfate is in the right test tube.

Here's another look at the barium sulfate precipitate. Again, the blue color of the solution has faded somewhat as a dilute solution of copper(II) chloride has been formed.

The blue solid is indeed copper(II) sulfate.

So we've solved it - step 4 in my initial reaction scheme produced a concentrated solution of copper(II) chloride and a precipitate of copper(II) sulfate. Apparently, copper sulfate is not soluble (or maybe slightly soluble) in a copper chloride solution.

What's strange is that the chloride appeared at all. Adding sodium hydroxide in step 3 should have neutralized any remaining hydrochloric acid and yielded NaCl, and washing the precipitate in step 4 should have carried away most or all of this salt. I would not expect that there would be enough chloride left over to form such a concentrated solution of copper chloride.

So in the end, I figured out what exactly was in my solution which was quite satisfying, but I'm still unsure of why this happened. I'll update these posts once I do some more research and determine the cause. Over this three part saga I learned a few things and had a lot of fun along the way, and I hope you did too.


  1. Could it be that you used glassware that wasn't completely clean of Cl contamination?
    The only other thing I could imagine is if there was some Cl anions in your H2SO4.

    It's been over a year, have you given any more thought as to what it could have been?

  2. My thoughts, when you precipitate copper (II) hydroxide, what is really happening is a ligand exchange on [Cu(H2O)6]2+ by two hydroxide ions to form the insoluble complex dihydroxytetraaquacopper(II) (which can be thought of as copper (II) hydroxide tetrahydrate). With such a high chloride conc. in the copper chloride solution, you really have green [Cu(H2O)2Cl4]2- anions floating around, which also undergo a ligand exchange with hydroxide to form what I would guess to be hydrated copper hydroxychloride (CuOHCl * 4H2O). This means your precipitate contained complexed chloride, and these disassociated when dissolved in H2SO4, where the dilute Cl- concentration promoted [Cu(H2O)6]2+ formation, giving blue color back.

  3. Hi, at the school we done this experiment 2 AgNO3(aq) + CuCl2(aq) then we add more AgNO3 to reach equilibrium so the concentration of the reaction increased the reaction shift to right to reach equilibrium the was a solid or precipitate. could you help me to answer this question?
    Does precipitation mean that the concentration of the compound that is the solid increases or decreases in solution?

  4. Hi,
    The precipitation does increase the concentration of reactants driving the formation of the products including the precipitate.
    But the concentration of a solid doesn't depend on the amount of solid, therefore the amount of precipitate increases but not its concentration.
    Also, since precipitation occurs after the solution is saturated, the concentration of the compound AgCl doesn't increase on adding more of reactants, the excess just precipitates.

    Hope this explains the previous query

  5. Heavy metal chlorides are bad for precipitation. You usually get a chloro complex with the insoluble group. If I need a heavy metal carbonate or hydroxide, I make sure I start with a nitrate or sulfate of the metal.