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EquationProblems1
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I was needing a little help balancing an equation: I'm exploring various ways of making NH4NO3. (And if your contribution is to tell me I'll end up on some watch list for even mentioning the compound, please keep your ridiculous ill-informed paranoia to yourself - if I was a terrorist, I'd be buying NH4NO3 fertiliser, not making a few grams!).

 

Having successfully mixed NaHSO4 with NaNO3 and neutralised with NH3, I balanced to give Na2SO4 + NH4NO3. So I replaced the nitrate salt with KNO3, and expected an eqimolar precipitate of Na2SO4 and K2SO4 (and of course the NH4NO3 in solution). However, I've tried a few equation balancers online, and they tell me it's an "impossible reaction". Am I wrong, or are they?

 

(My reasoning is that the NO3 and the NH3 combine, mopping up the extra H's liberated by the highly acidic NaHSO4/KNO3 mixture. This leaves K and 2Na and SO4 floating around combining with one another)


S'Ok - think I've done it:

 

2 NaHSO4 + 2 KNO3 + 2 NH3 = K2SO4 + Na2SO4 + 2 NH4NO3

Edited by EquationProblems1
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However, I've tried a few equation balancers online, and they tell me it's an "impossible reaction". Am I wrong, or are they?

 

Even if equation is balanced on "paper", it does not mean such reaction will happen in reality..

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On the contrary, it does mean it will because a reaction will occur based on the laws of empalthy. Please ensure contributions show actual knowledge of the subject matter.

Studio - I've done so, and KNO3 dissolves just fine. Much better 8n warm water though. But what''s your point?

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1)On the contrary, it does mean it will because a reaction will occur based on the laws of empalthy. Please ensure contributions show actual knowledge of the subject matter.

2)Studio - I've done so, and KNO3 dissolves just fine. Much better 8n warm water though. But what''s your point?

 

1) Enthalpy has nothing to do with it. (please note correct spelling of the term)

 

2) I suggested you look at the solubility of sodium and potassium sulphates, since your proposed reaction equation shows them precipitating. Yet you report the solubility of potassium nitric acid in water?

If you don't understand the significance of these solubilities you should not be fiddling with any quantities of these chemicals and I am not going to tell you.

Edited by studiot
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Jeez, I thought I was posting on a forum where people had at least a rudimentary knowledge of chemistry! I didn't think I'd need to point out that solubility varies with temperature, given that this is after all the whole basis of recrystallization. Think it might be you reaching for that textbook ...

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Jeez, I thought I was posting on a forum where people had at least a rudimentary knowledge of chemistry! I didn't think I'd need to point out that solubility varies with temperature, given that this is after all the whole basis of recrystallization. Think it might be you reaching for that textbook ...

Acting hostile towards people trying to point you in the direction of very simple chemistry is not the way to go here. Any person with a knowledge of fundamental chemistry would well know that salts containing first row metals are fully soluble in water (as a general rule). This is stark contrast with your claim that you would get, 'an eqimolar precipitate of Na2SO4 and K2SO4.'

Studiot, may I ask you not to post againn as you're not a chemist, nor a physicist if you're unaware that enthalpy drives the direction of a reaction.

Should also point out that enthalpy is only one factor that determines whether or not a reaction will occur spontaneously. You very much have entropy working against you in your equation.

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The simple fact is that you are using the wrong salts.

 

What you have not said is why this is in Homework Help ?

 

You have certainly complied with the special SF rules in this section in that you have shown an attempt.

 

Unfortunately you have also quite rudely resisted attempts to help you.

 

Remember that we do not do you homework for you, just help you complete it.

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EquationsProbles1

 

Yes, your equation is properly balanced. Your problem with this lies insolubility. Both of your sulfate products are very, very, soluble in water. What this means is that the products never actually form-- the individual ions stay in solution and you can only get the sulfates out by evaporating the water. As for the Ammonium nitrate, it is also soluble in water to the tune of about 20 moles per liter (118 grams per 100 ml according to my old handbook of chemistry and physics). When all the products of a potential reaction are soluble, the final compounds are never formed-- its just a big soup of ions floating around together (and fyi-- If I recall correctly, its not enthalpy that determines if a reaction if favorable, its Gibbs Free Energy, which also takes into account Entropy. All the enthalpy change won't do it if entropy is too unfavorable).

 

I realize I'm violating the homework rules by giving an answer, but this discussion seems to have dragged on.

Edited by OldChemE
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Thanks OldChemE. Appreciated.

 

I heated to about 65 Celsius for 75 minutes to reduce the solution by about 50%, then cooled to 2 Celcius overnight and filtered in the morning. A large amount of crystals had precipitated which I tested for nitrates and found none/very little to be present. So I'll warm the filtrate today to evaporate off all the water and see whether it yields the ammonium nitrate I'm hoping for. Will revert in a day or so.

 

If it does, I'll guess I'll need an assay for NH4NO3 :) - as you say, there are all sorts of ions floating about so it'll be interesting to see what has combined.

Edited by EquationProblems1
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You are clearly are interested in Chemistry, but you still haven't said why this is in homework?

 

I fully endorse what OldChemE said and you should look up the terms if you have not met them.

 

It is Free Energy (and in particluar Gibbs free energy as there are several types) that are the thermodynamic potentials that allow a reaction to happen spontaneously, not enthalpy alone.

Free energy is a combination of enthalpy and entropy.

 

Note I said spontaneously because it may well be possible to force a reaction by adding energy.

 

Further, the free energy says nothing about the rate of said reaction. Just because a reaction is thermodynamically possible, it does not mean it will go very far because the rate may be so slow as to be negligable.

For example silica is soluble in water, but over a timescale of tens of millions of years.

 

OK back to your particular quest.

 

There is nothing more satisfying than working something out for yourself. I had hoped the hint link to Wiki would cause you to realise the solubility of (ammonium) nitrate as you looked down the page and then you might also see the other information presented.

As noted, group 1 or the alkali metals groups compounds are generally very soluble, as are ammonium compounds.

But group 2 or the alkaline earths have many insoluble compounds.

This fact is often used in schemes such as yours to precipitate the unwanted ions.

 

I said you are using the wrong reagents.

So here is a scheme mixing a soluble calcium compound with a soluble ammonium compound to produce an insoluble calcium compound.

[math]2NH_4^ + \left( {aq} \right) + SO_4^{2 - }\left( {aq} \right) + C{a^{2 + }}\left( {aq} \right) + 2NO_3^ - \left( {aq} \right) \to CaS{O_4} \downarrow + 2\left( {NH_4^ + + NO_3^ - } \right)\left( {aq} \right)[/math]

Did you notice the way I wrote the equation, showing the precipitate?

Note also that charges must balance as well as atomic proportions when you right an equation like this.

And also that you need to find suitable soluble/insoluble pairs to make the scheme work?

 

 

 

Edit

 

A final word of caution since you are actually doing practical work.

although ammonium compounds in solution are pretty safe, Ammonium nitrate is not the only explosive ammonium compound.

Edited by studiot
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