sternman318 - 23-4-2011 at 12:58
So I was trying to make a solution of zinc by dropping some zinc from a battery into a copper sulphate solution. The copper solution was pretty
crappy, as I had messed around with it ( electolysis of it with a silver plated fork at one point haha), but it was pretty concentrated and it worked
great. The zinc I had used today had a bit of the MnO2 from the battery still stuck on it ( it was hard to remove and i just figured I would filter it
out with the copper in the end). So a lot of copper precipitated out, and eventually the zinc just sat and bubbled, so I assumed that all of the Cu2+
had reacted and that the zinc was reacting with the water. But-
1. Zn +2H2O -> Zn(OH)2 + H2 , right? Wouldn't the Zn(OH)2 prevent further reaction? The metal continued to bubble, and appeared black. One
explananation is that the solution could have been acidic from the electrolysis , but I doubt it ( it didnt work too well)
2. The solution, mopped up by the zinc, should have been colorless( free of Cu2+). But, it is a pale green- could this be from the manganese? Adding
Zn or Mg do not yield a copper precipitate, so I think the solution is free of Cu2+
[Edited on 23-4-2011 by sternman318]
sternman318 - 23-4-2011 at 18:01
I know color isn't really good indicator as it can vary so much, but here is a picture
blogfast25 - 24-4-2011 at 04:30
You mention 'manganate' and that is indeed the only type of manganese ion that would be green: MnO4 (2-) with Mn in the oxidation state +VI.
Manganate is the precursor to permanganate: MnO4 (-) (oxidation state +VII). Neither however can arise in the conditions you
describe.
The green is likely to be residual Cu2+, with some chloride it forms green CuCl4(2-) anions...