Doyle3694
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Prediction of chemical reaction
Hello!
I have a passionate interest in chemistry and understand many concepts quite well, but it feels like all I ever get to is remembering what reacts and
what doesn't, without any knowledge of rule about what would react why.
What I am interested in here is mostly the double displacement reactions, how do I predict if 2 salts are going to swap anions or not when dissolved
in water? What I have been looking at so far is both reactivity and electronegativity which I guess play a role in these kinds of reactions? Problem
is, most lists and so on of these don't cover the anions which consists of more than 1 atom, like hydroxides or nitrates. Similarily they don't cover
for example ammonia in relation to the other cations.
In practical sense, I was thinking about the synthesis of anhydrous nitrate from the calcium ammonium nitrate I can find in the coldpacks that are
sold locally.
Problem is, I don't have access to the traditional potassium chloride which is what is usually used here and instead theorized using sodium
bicarbonate and then heating, drying, redissolving, filtering and drying again to get out the sodium nitrate. In this case you would have the overall
reaction be
5Ca(NO3)2 * NH4NO3 * 10H20 + 11NaHCO3 -> 11NaNO3 + 5Ca(HCO3)2 + NH4HCO3 + 10H2O
(which way it goes here is pretty unclear, CAN seems to be either 5 calcium 1 ammonia 10 water or 5 ammonia 1 calcium 10 water)
However, I can't find any source if this would work or not. What is the sort of "wikipedia stat" that would help predict what anion will go where? Is
there such a stat?
TL/DR What decides which anions go where in a double displacement reaction?
[Edited on 2-1-2016 by Doyle3694]
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byko3y
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Real world decides which anion goes where.
WTF is "anhydrous nitrate"? Nitrate of some random metal? Then you can just heat your salt until decomposition of ammonium nitrate is finished.
Calcium nitrate decomposition starts somewhere at 500°C.
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Doyle3694
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Yeah just want my nitrate to be without any water simply, so sodium or potassium nitrate are prime targets. Calcium nitrate tetrahydrate boils at 132
degrees C which is lower than the decomposition temperature(And it is also fairly toxic). And as I said before, what I have access to is a doublesalt
consisting of both calcium nitrate and ammonium nitrate.
What I am trying to get at is are there any rules for what takes what anion at all? Why would potassium chloride always swap anion with ammonium
nitrate to form potassium nitrate and ammonium chloride?
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byko3y
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132 C is a boiling point of hydrate, obviously. Anhydrous calcium nitrate has extremely high boiling point.
Quote: Originally posted by Doyle3694  | | Why would potassium chloride always swap anion with ammonium nitrate to form potassium nitrate and ammonium chloride? |
Because those salts are fully dissociated in a solution, so they will randomly assemble on crystallization into one of the 4
possible compounds.
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Doyle3694
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Then why when this solution is cooled does potassium nitrate at a fairly high purity crystallize? Why is the end result not heavily contaminated with
potassium chloride?
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blogfast25
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| Quote: Originally posted by Doyle3694 |
I have a passionate interest in chemistry and understand many concepts quite well, but it feels like all I ever get to is remembering what reacts and
what doesn't, without any knowledge of rule about what would react why.
What I am interested in here is mostly the double displacement reactions, how do I predict if 2 salts are going to swap anions or not when dissolved
in water? What I have been looking at so far is both reactivity and electronegativity which I guess play a role in these kinds of reactions? Problem
is, most lists and so on of these don't cover the anions which consists of more than 1 atom, like hydroxides or nitrates. Similarily they don't cover
for example ammonia in relation to the other cations.
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These are such broad questions they're almost impossible to answer.
Specifically in the case of displacement reactions it mostly boils down to solubility, to determine what will precipitate or crystallise out.
In the case of the method you propose to recover soluble nitrates from cold packs, it would be better to use Na2CO3 than
NaHCO3 because CaCO3 is very insoluble but Ca(HCO3)2 is not.
5 Ca(NO3)2.NH4NO3.10H2O(aq) + 5 Na2CO3(aq) -> 10 NaNO3(aq) + 5 CaCO3(s) + NH4NO3(aq) + 10 H2O(l)
After filtering out the CaCO3 you'd be left with a solution of NaNO3 and NH4NO3.
Wikipedia's solubility table:
https://en.wikipedia.org/wiki/Solubility_table#A
... tells us that the latter is far more soluble than the former, which should allow partial separation by thermal crystallisation: boil off a
calculated amount of water and on cooling you should get crystals of NaNO3.
[Edited on 2-1-2016 by blogfast25]
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MolecularWorld
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Or add sufficient sodium carbonate to react with both the calcium nitrate and the ammonium nitrate, filter the calcium carbonate as above, and boil to
dryness. The boiling should drive off the ammonia and associated carbonate.
Edit: Incidentally, I recently used a method similar to this to make purified calcium nitrate from cold pack calcium ammonium nitrate. By distilling a
solution of the calcium ammonium nitrate with excess calcium hydroxide (in my case, conveniently available as "pickling lime" a few aisles over from
the cold packs), I was able to recover most of the ammonia as a solution. Decanting the hot liquid left in the stillpot from the undissolved calcium
hydroxide gave a solution of calcium nitrate with slight calcium hydroxide contamination, which was removed via the addition of carbonated water,
which precipitated the calcium hydroxide as carbonate.
[Edited on 2-1-2016 by MolecularWorld]
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Doyle3694
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My initial thought was to use heat to decompose the ammonium bicarbonate, and I thought it would be more trouble to decompose ammonium carbonate into
ammonium bicarbonate and then decompose it again, releasing double the amount of ammonia, than it would be to dry it so the calcium bicarbonate
becomes calcium carbonate and then redissolve it filtering out the calcium carbonate.
Thanks for the help! Sadly as with many things in science this just leaves me more confused than when I started. Are there any rules at all or is it
simply a case to case between what salts are involved? And if so, is the only way to know what would happen in a proposed displacement reaction to
actually test the reaction in the lab?
If you know any good literature on the subject I would love some links!
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Doyle3694
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| Quote: Originally posted by MolecularWorld | | Or add sufficient sodium carbonate to react with both the calcium nitrate and the ammonium nitrate, filter the calcium carbonate as above, and boil to
dryness. The boiling should drive off the ammonia and associated carbonate. |
This was exactly my theorized method of getting rid of the ammonia
[Edited on 2-1-2016 by Doyle3694]
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blogfast25
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| Quote: Originally posted by MolecularWorld | | Or add sufficient sodium carbonate to react with both the calcium nitrate and the ammonium nitrate, filter the calcium carbonate as above, and boil to
dryness. The boiling should drive off the ammonia and associated carbonate. |
Yup, good idea.
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blogfast25
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| Quote: Originally posted by Doyle3694 | Are there any rules at all or is it simply a case to case between what salts are involved? And if so, is the only way to know what would happen in a
proposed displacement reaction to actually test the reaction in the lab?
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For reactions in general there's a whole body of theory that helps predict what is possible and what's not. Far too large to summarise here.
But for these aqueous displacement reactions solubilities are the main predictor. If a possible anion/cation combination is insoluble or poorly
soluble then that will drive the reaction.
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byko3y
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| Quote: Originally posted by Doyle3694 | | Then why when this solution is cooled does potassium nitrate at a fairly high purity crystallize? Why is the end result not heavily contaminated with
potassium chloride? |
At 0°C KNO3 solubility is two times lower than that of KCl and NH4Cl.
Fast cooling of a strongly saturated solution will lead to KCl and NH4Cl contamination.
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Doyle3694
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| Quote: Originally posted by blogfast25 |
But for these aqueous displacement reactions solubilities are the main predictor. If a possible anion/cation combination is insoluble or poorly
soluble then that will drive the reaction. |
Does this also mean that discarding any thermodynamic effects the possible salts will generally fall out of solution upon removal of water or changed
temperature in the order of their solubility in water?
[Edited on 2-1-2016 by Doyle3694]
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Doyle3694
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Are the boiling temperatures of hydrates always specified as when they start releasing their water and not as when they themselves boil? If I
understood you correctly, do all hydrates boil at 132 C?
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DraconicAcid
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| Quote: Originally posted by Doyle3694 |
Are the boiling temperatures of hydrates always specified as when they start releasing their water and not as when they themselves boil? If I
understood you correctly, do all hydrates boil at 132 C? |
Most sources will give a decomposition temperature for a hydrate rather than a boiling point, because ionic compounds do not generally boil at
temperatures below 600 oC. And these will be at at a range of different temperatures, not all 132 oC.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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Doyle3694
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| Quote: Originally posted by DraconicAcid |
Most sources will give a decomposition temperature for a hydrate rather than a boiling point, because ionic compounds do not generally boil at
temperatures below 600 oC. And these will be at at a range of different temperatures, not all 132 oC. |
This would also mean that heating Ca(NO3)2 * 4H2O would only release water(Before the Ca(NO3)2
itself decomposes ofcourse) and what wikipedia deems as it's boiling point is just the point where it loses it's water of crystallization?
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blogfast25
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| Quote: Originally posted by Doyle3694 |
This would also mean that heating Ca(NO3)2 * 4H2O would only release water(Before the Ca(NO3)2
itself decomposes ofcourse) and what wikipedia deems as it's boiling point is just the point where it loses it's water of crystallization?
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It's not a question of 'deeming'. Assume a hydrate melts at a certain temperature and then on further heating that melt starts boiling then that is
the BP (say at STP) of that substance. It happens also to be the temperature where it starts losing water quickly or in other cases where it starts
hydrolysing fast.
The BP of a hydrate is a real BP, that's why Wiki calls it that.
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Doyle3694
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Instead of getting triggered over my choice of words, can you answer the question?
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MrHomeScientist
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As others have alluded to, the result of a potential double displacement reaction (or in general really any aqueous chemistry) is going to depend on
solubility of the possible products. The more modern way to think of things is as a mixture of ions.
When you mix, for example, sodium nitrate and potassium chloride, a reaction you might expect is: NaNO<sub>3</sub> + KCl ==
KNO<sub>3</sub> + NaCl
Will this actually happen? First off, it's important to note that in this case, when you mix solutions of the two reactants together nothing actually
happens. There is no 'reaction' that yields KNO<sub>3</sub> molecules in solution. Rather, you simply have a mixture of ions in solution:
K<sup>+</sup>, NO<sub>3</sub><sup>-</sup>, Na<sup>+</sup>, and Cl<sup>-</sup>. You then
have to determine the solubility of each possible combination of ions. Using Wikipedia's solubility table, we find (at 10C, in units of g/100g water):
NaNO<sub>3</sub> - 80.8
KCl - 31.2
KNO<sub>3</sub> - 22
NaCl - 35.72
From this you can see that potassium nitrate is the least soluble, and therefore will crystallize first when the solution is cooled. So you could say
that yes indeed, a 'reaction' has occurred and you have produced new products.
Even if there is some precipitation of chlorides, they should be easy to spot since they are cubic crystals as opposed to the nitrate's long thin
needles. If purity is a concern, harvest a crop of crystals, redissolve them in water, and evaporate the solution to ~10% remaining liquid. The
impurities will be concentrated in the liquid, and can be discarded. One or two recrystallizations will generally yield a very highly pure substance.
Also from the solubilities you can see that sodium nitrate is extremely soluble even at low temperatures, thus it is quite a bit harder to isolate and
purify. I'd recommend making potassium nitrate instead if the cation doesn't matter. It has a great solubility curve, being 10x more soluble in
boiling water than in freezing water, so it is very easy to get a lot of pure product. Dissolve as much KNO<sub>3</sub> in boiling water
as you can, then cool to near freezing and lots of beautiful needles crystallize out.
The suggestion to use sodium carbonate (washing soda) is excellent, too. Combining that with your calcium ammonium nitrate will yield calcium
carbonate (easily filtered off) and ammonium carbonate (which is unstable and decomposes on heating into only gaseous products). This leaves only
sodium nitrate left in solution, which can be evaporated to dryness.
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Doyle3694
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Thanks MrHomeScientist, just the answer I was looking for
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Praxichys
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| Quote: Originally posted by MrHomeScientist | (at 10C, in units of g/100g water):
NaNO<sub>3</sub> - 80.8
KCl - 31.2
KNO<sub>3</sub> - 22
NaCl - 35.72
From this you can see that potassium nitrate is the least soluble, and therefore will crystallize first when the solution is cooled.
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Technically yes, but in practice these solubilities are far too close to practically isolate anything
economically. Note that (excluding the common ion effect) KNO3 and KCl differ by only 9.2g/100ml @ 20C. This means that out of 22 possible grams of
recoverable KNO3, only 9.2 can be isolated. Re-boiling to further concentrate followed by cooling again will not work since there will be a nitrate
deficiency from the last precipitation, and it will fall out of solution mixed with everything else. In reality, the common-ion effect of the highly
soluble sodium nitrate will tend to force NaCl out much earlier than the solubility data suggests, also contaminating the initial precipitation of
KNO3.
A bit better of a method would be to boil the solution down until a substantial amount of the chlorides had precipitated, then filter while hot to end
up with a solution of the more soluble components, the KNO3/NaNO3. Either way, it would be a contaminated mess.
| Quote: Originally posted by MrHomeScientist | | The suggestion to use sodium carbonate (washing soda) is excellent, too. Combining that with your calcium ammonium nitrate will yield calcium
carbonate (easily filtered off) and ammonium carbonate (which is unstable and decomposes on heating into only gaseous products). This leaves only
sodium nitrate left in solution, which can be evaporated to dryness. |
I second this as a great solution. It
will need prolonged boiling to drive off the ammonium carbonate though.
I have been successful manufacturing ammonium nitrate from CAN fertilizer by adding ammonium sulfate, which is another cheap fertilizer. The two are
dissolved in hot water and poured together. Calcium sulfate precipitates immediately and almost completely and can be filtered en masse through fabric
(I used a t-shirt) over a bucket. The resulting solution is boiled down and left to cool, and the ammonium nitrate crystallizes from the syrup, which
is then baked dry and pulverized. I know you're looking for Na/K salts, but if you ever need a good route to ammonium nitrate, this works in a similar
way. Notice that this is a solubility-driven reaction, just like the carbonate method mentioned above.
If you are trying to get to 100% nitric acid, distilling ammonium nitrate with sulfuric acid works but some water is formed as a by-product and the
product will not be pure nitric acid. If your goal is to run a nitration using sulfuric acid and nitrate salts, ammonium nitrate is ideal since the
metal sulfates that are present when using Na or K nitrates tend to precipitate at low temperatures and make the reaction difficult to stir.
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MolecularWorld
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| Quote: Originally posted by Praxichys | | I have been successful manufacturing ammonium nitrate from CAN fertilizer by adding ammonium sulfate, which is another cheap fertilizer. The two are
dissolved in hot water and poured together. Calcium sulfate precipitates immediately and almost completely and can be filtered en masse through fabric
(I used a t-shirt) over a bucket. The resulting solution is boiled down and left to cool, and the ammonium nitrate crystallizes from the syrup, which
is then baked dry and pulverized. I know you're looking for Na/K salts, but if you ever need a good route to ammonium nitrate, this works in a similar
way. Notice that this is a solubility-driven reaction, just like the carbonate method mentioned above. |
This definitely works, though as you know, the calcium sulfate contamination may be greater than expected. It may be possible to remove the calcium sulfate impurity with methanol, though I haven't tried this and can't find any good data on the solubility of calcium sulfate in methanol.
| Quote: | | If you are trying to get to 100% nitric acid, distilling ammonium nitrate with sulfuric acid works but some water is formed as a by-product and the
product will not be pure nitric acid. |
Have you done this? Did you use a vacuum? I'm led to believe (by, for
example, one of the oldest extant threads) that there's a risk of runaway decomposition, especially if this is done at atmospheric pressure.
[Edited on 8-1-2016 by MolecularWorld]
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Praxichys
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| Quote: Originally posted by MolecularWorld |
| Quote: | | If you are trying to get to 100% nitric acid, distilling ammonium nitrate with sulfuric acid works but some water is formed as a by-product and the
product will not be pure nitric acid. |
Have you done this? Did you use a vacuum? I'm led to believe (by, for
example, one of the oldest extant threads) that there's a risk of runaway decomposition, especially if this is done at atmospheric pressure.
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Yes, many times, and never under reduced pressure - that was my primary means of obtaining nitric acid in
the early years of the lab. The procedure involved 2eq of sulfuric acid to one of ammonium nitrate, mainly to keep the ammonium sulfate from forming a
solid brick at the end, so I could do repeated runs rather rapidly. I used prills from cold packs and 93% drain opener sulfuric acid, in a 500ml RBF.
After the first distillation, the nearly pure and NO2-contaminated acid was then diluted with tap water and redistilled to replenish my stock of 68%
azeotropic acid. Invariably there was a small amount of ammonium nitrate left in the flask after the second distillation - I had always attributed
this to the volatility of ammonium salts but later found that some of the ammonium sulfate was probably decomposing to ammonia and ammonium bisulfate.
I never experienced anything resembling thermal runaway. Nitric acid volatilizes at 83C, so I really doubt there is any substantial amount of nitrate
left in the flask by the time it gets to a temperature unreasonable for ammonium nitrate to be at.
But, take this advise with a grain of salt - I was working with small batch sizes on a gentle heating device, using twice as much 93% acid as
technically required. Perhaps I was very close to disaster and didn't know it.
[Edited on 8-1-2016 by Praxichys]
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