I am confused when it comes to predicting the products of a reaction when two inorganic salts/bases/acids in solution are mixed.
For instance, I understand that:
NaNO3 + KCl ---> KNO3 + NaCl
I assume this is because: (A) potassium cations are more electropositive than sodium cations, and (B) nitrate anions are more electronegative than
chloride anions, thus (C) the potassium and nitrate ions will have the strongest affinity for each other.
So, am I missing anything here? At this point, the HSAB theory is a bit over my head.
The part I get confused on is when it comes to hydroxide and ammonium ions. For example:
Will this work?
2NH4NO3 + K2SO4 ---> (NH4)2SO4 + 2KNO3
And this?
Ca(OH)2 + 2NH4NO3 ---> Ca(NO3)2 + 2NH4OH
In the first reaction (if feasible), is the ammonium cation and the sulfate anion the most active, thus having the strongest affinity for each other?
What about the second reaction; will it take place, and why?Pyridinium - 23-3-2007 at 07:17
In the sense of a reaction involving two aqueous salts, they don't react in solution, because both remain fully ionized. Therefore, as long as both
remain dissolved, there technically hasn't been any reaction yet.
The change happens when it's time to crystallize one or more solutes.
In your example where K+ and NO3- have the most affinity for each other, you'll also notice that KNO3 has the lowest solubility of the four possible
salts. It will crystallize first out of solution.
EDIT: where I said above, "they don't react in solution", that assumes you're not dealing with something that immediately precipitates. (such as,
Ca++ ions and SO4-- ions)
[Edited on 23-3-2007 by Pyridinium]obsessed_chemist - 23-3-2007 at 07:52
Recently I tried making saltpeter by mixing solutions of potassium sulfate with ammonium nitrate. The potassium sulfate was made with KCl and epsom
salts, Then easily seperated through fractional crystallization.
For some reason though, after mixing the ammonium nitrate and potassium sulfate solutions, boiling down, chilling, and fractionally crystallizing, the
potassium nitrate that precipitated was fairly impure. My thought is that the ammonium sulfate by-product decomposed into ammonia and sulfuric acid
while boiling, and the residual sulfuric perhaps reacted with the potassium nitrate to form potassium sulfate, which is even less soluble. Next time I
plan on using a pH indicator, and frequently adding household ammonia to the boiling solution to keep the pH from dropping.
I really want to get this procedure right, because it could be a really cool OTC way to get saltpeter. The only ingredients needed would be:
-Epsom salt
-KCl salt substitute, or water-softener
-Instant cold-packs
-Clear Ammonia
I recently discovered that potassium sulfate is available at any garden center, but then again, you could just buy KNO3 stump remover or NaNO3, etc
while you're there. My intention is to improvise a completely OTC saltpeter prep from things available at Wal*Mart.Pyridinium - 23-3-2007 at 08:09
Quote:
Originally posted by obsessed_chemist
For some reason though, after mixing the ammonium nitrate and potassium sulfate solutions, boiling down, chilling, and fractionally crystallizing, the
potassium nitrate that precipitated was fairly impure.
You might get a better product by doing extremely slow evaporation using a seed crystal of KNO3, just make sure the solution isn't so dilute that it
eats up your seed crystal.
Fast crystallization or precipitation almost always traps impurities.obsessed_chemist - 23-3-2007 at 08:19
Quote:
Originally posted by PyridiniumYou might get a better product by doing extremely slow evaporation using a seed crystal of KNO3, just make
sure the solution isn't so dilute that it eats up your seed crystal.
Fast crystallization or precipitation almost always traps impurities.
Thanks for confirming that. I had a feeling that was the case. I definitley chilled the solution in a hurry.
Any idea about the ammonium sulfate decomposing while boiling?not_important - 23-3-2007 at 09:41
Ammonium sulfate does lose ammonia when a solution of it is boiled, the acid sulfate is formed (NH4)HSO4 and the loss of ammonia is not complete so
you have a mix of normal and acid sulfates.
Try to use as close to saturated solutions to start with, to reduce the amount of evaporation needed.
Potassium sulfate has a fairly low solubility, using the KCl might be better. Try graphing the solubility curves, don't plot weight but rather moles
and take into account relationships like 2 KCl <=> K2SO4 when looking at the plots. Once it looks as if you have a combination that will work,
then convert back to the weight needed.
The other thing to remember is that sometimes you have a case where the original mixing (hot) will result in a portion of one product crystallising
out, and cooling (after filtering) will cause the other product to drop out. When that happens you usually need to recrystallise the 2nd product to
clean it up as some ofthe 1st product will be joining it.obsessed_chemist - 23-3-2007 at 10:00
For some reason I was under the impression that ammonium nitrate and potassium chloride couldn't be combined to make KNO3. But since metathesis is
apparently soley based on solubility, and not reactivity of ions, I guess I'm gonna have to rework a lot of things.
Also, the solubility of KCl and KNO3 are way too close to each other for this to work.
However, mixing potassium hydroxide solution (widely available here as drain cleaner in 45% solution) with NH4NO3 solution, then cooling may work.
This might have to be done in a cold phase with much control to prevent ammonia from bellowing out of the reaction vessel. Since KNO3 is way less
soluble than either KOH, NH4NO3, or NH3/NH4OH, this method should work really well. Not quite as OTC as the afforementioned, but viable
none-the-less.
[Edited on 3/23/2007 by obsessed_chemist]12AX7 - 23-3-2007 at 10:24
If you don't need to save the NH3, you can cook it off just as well.
TimMagpie - 23-3-2007 at 15:03
I think don't one can make general rules about the outcome of theorized methathesis reactions. Evolution of a gas or formation of a precipitate can
drive to a recoverable reaction product. But one can also form double salts, etc, under the right conditions of concentration and temperature. I
think you have to look at each reaction situation in detail. And then, assuming you don't have access to predictive software, you may have to
determine the outcome by experiment.
I also agree that while all the ions are still in solution you really don't have a reaction, ie, the cations are not "owned" by certain anions and
vice versa.Levi - 23-3-2007 at 16:09
Quote:
Originally posted by Magpie
I also agree that while all the ions are still in solution you really don't have a reaction, ie, the cations are not "owned" by certain anions and
vice versa.
This is an extremely important concept that isn't hammered home like it should be in general chem courses. That said, I believe that some ions do
associate more closely with one another in solution. This isn't very important since they can't be recovered in such a state but I'm pretty sure
there is a related temperature change that isn't accounted for by energy of solvation. Am I just blowing smoke or is this correct?Magpie - 23-3-2007 at 16:24
Levi, you have a point that something is going on when there are heats of solution, density changes, etc, upon mixing. Hydrogen bonding and soluble
coordination complexes come immediately to mind.Pyridinium - 24-3-2007 at 00:25
Quote:
Originally posted by Magpie
Levi, you have a point that something is going on when there are heats of solution, density changes, etc, upon mixing. Hydrogen bonding and soluble
coordination complexes come immediately to mind.
This is true. However, while it may be a bit of an oversimplification to consider ions unreacted until the first moment of crystallization, it works
for general purposes. As long as there are no precipitates or crystals, I consider any dissolved complexes (H-bonded, etc) as reversible unless there
is some way to detect them, such as color change.
I suspect a great deal of work has been done in this area, with some fine articles buried way back in the literature. As Magpie pointed out, things
can get pretty complex with double salts, co-crystallization, etc. Crystallizing something can be a very touchy process that is very
condition-dependent.obsessed_chemist - 29-3-2007 at 11:32
How can the fact that, aqueous solutions of ammonium hydroxide and calcium chloride don't precipitate calcium hydroxide upon mixing, be explained?Magpie - 29-3-2007 at 11:47
What is the dissociation constant of NH4OH? What is the Ksp of Ca(OH)2? Check the math, it may tell you why there is no ppt.obsessed_chemist - 29-3-2007 at 12:10
^ I obviously am in over-my-head with some of this stuff; not because of an inability to grasp, but merely a lack of academic knowledge on my part.
I shall read about "dissociation constant" theory, and report back.
