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Topic: dissolving ammonium nitrate in water  (Read 52771 times)

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Offline kevinkevin

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dissolving ammonium nitrate in water
« on: August 23, 2012, 03:57:05 PM »
  I just read a chemical safety sheet about ammonium nitrate and it said that when it is dissolved in water it makes the solution slightly acidic.  I wrote out some balanced equations, in they end they somewhat contradict each other, so I would like to tell you my conclusion on why I believe it makes the solution slightly acidic.

 First ammonium nitrate is dissolved in water; NH4NO3 + H2O --> NH4+ + NO3-

  Then I thought that this reaction could occur; NH4+ + H2O <--> NH3 + H3O+
  That would make the solution slightly acidic but now ammonia is present so this reaction would occur;
  NH3 + H2O <--> NH4+ + OH-
 That would favor the solution being slightly basic. 
  Now the only reason I could think of to explain why it forms a slightly acidic solution is because the first reaction that produces hydronium ions has a larger K value than the reaction producing hydroxide ions.  If the k values were the same then the hydronium ions and hydroxide ions would just form water and the solution would remain neutral. 

  Just on a side note, I all ways thought that the solution was basic because I have gotten some on my hands and it feels pretty slimy. 

 Thanks for any help.
 

Offline Arkcon

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Re: dissolving ammonium nitrate in water
« Reply #1 on: August 23, 2012, 05:06:42 PM »
There is another counter ion that you're ignoring.
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Offline kevinkevin

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Re: dissolving ammonium nitrate in water
« Reply #2 on: August 24, 2012, 01:30:58 PM »
  Is it the nitrate ion? Does it behave like this?...   NO3- + H2O <---> HNO3 + OH-    It seems to me like that reaction would not occur, but I am a first year student =(, if it does occur then the nitric acid should dissociate like so; HNO3 + H2O<--->NO3- +H3O+
So once again there is another set of equations where one produces hydroxide ions and the other produces hydronium ions.  So is the solution is slightly acidic the K value for the second equation must be higher than the first.  Is this on the right track or have i gone down the wrong path?   


   On a side note since the nitrate ion is the conjugate base of nitric acid is this the only reaction that occurs involving the nitrate ion?    HNO3- + H2O <--> NO3- + H3O+  But the hydrogen that makes the nitric acid has to come from somewhere, that's why I think the first reaction that yields hydroxide ions has to occur.

Offline Rutherford

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Re: dissolving ammonium nitrate in water
« Reply #3 on: August 24, 2012, 02:49:15 PM »
Ka of NH4+ is 5.7*10-10 meaning that a small amount of NH3 will be made (so only a very small amount of OH- will be made), and a big amount of NH4+ will remain in the solution. The reaction of NH3 hydrolysis you wrote has NH4+ ions on the right side, so by the Le Chatelier's principle the equilibrium will be shifted to the left, so this means that OH- ions will be present only in traces and they will react with some of the produced H+, therefore some of the H+ left in the solution making it acidic.
(This is not so statically as I wrote, concentrations are changing all the time)

Offline Arkcon

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Re: dissolving ammonium nitrate in water
« Reply #4 on: August 24, 2012, 03:45:25 PM »
So, kevinkevin:, you have two choices for each of the two ions -- acid or base for each.  Now, the old saw goes, a salt of a weak base and a strong acid give a weakly acidic solution.  Now, can you use the Ka's of each ion, like Raderford: said, to determine what each is, a strong acid, and a weak base?  I'd heard that nitric acid is a strong acid and ammonium hydroxide is a weak base, but since you've come to these problems right about now, it would be better if you can use the Ka to prove it.
Hey, I'm not judging.  I just like to shoot straight.  I'm a man of science.

Offline kevinkevin

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Re: dissolving ammonium nitrate in water
« Reply #5 on: August 25, 2012, 02:42:30 PM »
  Sorry, ignore that post above me.
  I got a tiny bit lost on that last part.  Since NH4+ has such a small Ka value, very small amounts of OH- will be produced, but a large amount of NH4+ will remain in solution.  Will enough of the NH4+ react to form H3O+?  I also have one more concern.  Since NO3- is the conjugate base of a strong acid that will make it a very weak base correct, all most like a spectator ion with respect to making the solution acidic or basic.  On the other hand though, NH4+ is the conjugate acid of a weak base so wont that make it a stronger acid, even though still very weak?  So now to think of K values that would make sense.  Since the NH4+ reaction yields a hydronium ion and ammonia, ammonia's Kb value must be smaller than the Ka value NH4+.  Also for NO3-, which yields HNO3- and OH-, hardly ever occurs since it is the conjugate base of a strong acid.
  Thank-you for that information about salts Arkcon.=) Could you please explain the last part of your post to me though please?  Are you saying that I should find the K values for ammonium hydroxide and nitric acid?                   

Offline Rutherford

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Re: dissolving ammonium nitrate in water
« Reply #6 on: August 25, 2012, 03:18:28 PM »
NH3 is a weak base, but still not too weak ,Kb=1.8*10-5, the Ka of its conjugate acid can be calculated by the formula Ka*Kb=Kw, and you get a similar value I posted. Ka of HNO3 is 24, Kb of the conjugated base can be calculated by the previous equation. You will see that it is very very weak.

"Since NH4+ has such a small Ka value, very small amounts of OH- will be produced, but a large amount of NH4+ will remain in solution.  Will enough of the NH4+ react to form H3O+?"

NH4+ produces by hydrolysis only H3O+. NH3 produces OH-.

Offline kevinkevin

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Re: dissolving ammonium nitrate in water
« Reply #7 on: August 25, 2012, 05:27:16 PM »
  I am still confused by the balanced equations and the whole NH3 and NH4+ thing.  I think I might have gotten a little lost with my thoughts.  This is now specifically what I am confused about.   

  NH4+ + H2O <--> NH3 + H3O+   So this is how the hydronium ions are produced to make the solution slightly acidic correct?  Then the NH3 will produce the hydroxide ions.  But we don't want hydroxide ions we want hydronium ions.  Also we don't have HNO3 to make the solution more acidic, we just have NO3-, which seems to me will not effect the pH of the solution. 

  From what you guys have told me, if I interpreted it correctly, the reason the solution is slightly acidic is because the NH4+ will produce hydronium ions and NH3.  Then NH3 forms very little OH- because the concentration of ammonia is very small because of the very small K value of NH4+, and the K value for ammonia will not change of course so it has to work with the concentration of ammonia given from the NH4+ and H2O reaction.  So the result of all this is that more hydronium ions are left in solution because the ammonia can only be created from the NH4+ and H2O reaction, which yields very small amounts of NH3.  Correct?           

   

   

Offline Borek

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Re: dissolving ammonium nitrate in water
« Reply #8 on: August 25, 2012, 06:00:57 PM »
From what you guys have told me, if I interpreted it correctly, the reason the solution is slightly acidic is because the NH4+ will produce hydronium ions and NH3.  Then NH3 forms very little OH- because the concentration of ammonia is very small because of the very small K value of NH4+, and the K value for ammonia will not change of course so it has to work with the concentration of ammonia given from the NH4+ and H2O reaction.  So the result of all this is that more hydronium ions are left in solution because the ammonia can only be created from the NH4+ and H2O reaction, which yields very small amounts of NH3.  Correct?

Yep. You have NH4+ producing some small amounts of H+ and NH3, this small amount of ammonia produces even smaller amount of OH- - in the end concentration of OH- produced is much smaller than the amount of H+ produced.
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Offline kevinkevin

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Re: dissolving ammonium nitrate in water
« Reply #9 on: August 25, 2012, 09:25:58 PM »
   Exelent, nothing feels better than when a question is answered and you really understand it.   Thank-you guys!

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