Bill Stock wrote:

Thanks Rocco, so my MgSO4+7H2O should not reduce the solubility of my KCL if
I understand you correctly. This does not seem to be what I'm experiencing
though, as I'm achieving no where near the solubility numbers in Chuck's
calculator. Perhaps there is some CL in the Plantex mixture, which is
causing the KCL to precipitate.

Hmm.. I probably mislead you. In addition to each individual salt, you
also have to look at the double displacement salts (where you swap the
anions and cations from the two mixed salts).

The approximate solubility of various compounds in cold water, according
to the CRC handbook:

KCl - 350 g/L or 4.7 M for a Ksp of 22 (4.7*4.7)
MgSO4 - 260 g/L or 2.2 M for a Ksp of 4.8 (2.2*2.2)
MgCl2 - 543 g/L or 5.7 M for a Ksp of 740 (5.7*11.4*11.4)
K2SO4 - 120 g/L or 0.7 M for a Ksp of 1.4 (1.4*1.4*.7)

A saturated solution of KCl thus has 4.7 M K+ and 4.7 M Cl-
A saturated solution of MgSO4 thus has 2.2 M Mg++ and 2.2 M SO4--

A hypothetical "dual saturated" solution would have
4.7 M K+, 4.7 M Cl-, 2.2 M Mg++ and 2.2 M SO4--

But we also have to consider the double displacement products MgCl2 and
K2SO4.

The solubility products for the hypothetical solution are

MgCl2 - [Mg++][Cl-]^2 = 2.2*4.7*4.7 = 48.6, well under the 740 limit
K2SO4 - [K+]^2[SO4] = 4.7*4.7*2.2 = 48.6 - well *above* the limit of 1.4

So what's happening is that you dissolve the MgSO4 and KCl, and the two
dissolve fine, but when the K+ and the SO4-- find each other in
solution, they feel crowded and crash out as a precipitate of K2SO4.
That whitish "undissolved" powder in the bottom of the bottle is not KCl
or MgSO4, but K2SO4.

Note that you can probably get the precipitate to dissolve if you heat
it (The CRC handbook gives solubility of K2SO4 in hot water of twice
that in cold water), but when the solution cools the K2SO4 will just
come out of solution again.