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Topic: precipitates re-dissolve in nitric acid  (Read 17831 times)

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Offline breanainneire

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precipitates re-dissolve in nitric acid
« on: January 13, 2007, 01:37:11 PM »
Hi. My Chem 12 class just did a qualitative analysis lab (adding various solutions to other solutions to see if they would precipitate). In one case, I added Ba(NO3)2 to a solutions of Na2SO4 and Na2CO3 and precipitates formed. In another, AgNO3 was added to solutions of Na2CO3 and NaCl and also formed precipitates. We were told to add nitric acid to the solutions that formed precipitates. The BaCO3 dissolved completely, and the BaSO4, Ag2CO3 and AgCl dissolved a bit. What I’m wondering is is why do the precipitates redissolve? What reaction occurs (I’m looking for some kind of equation)? My teacher didn’t explain why and neither did the text. I’m kind of stumped.
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Offline xiankai

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Re: precipitates re-dissolve in nitric acid
« Reply #1 on: January 13, 2007, 10:52:39 PM »
i think it is called the common ion effect.

let's use the case of  Ba(NO3)2 and Na2SO4

say that all the ions are at equilibrium, and u add an excess of one ion (nitrate in this case). by le chatelier's principle, the system will act to reverse the effects of the excess. therefore take the following equation:

Ba2+ (aq) + 2NO3- (aq) --> Ba(NO3)2 (aq)

(or replace Ba2+ with Na+, but that isnt the reaction we're concerned about)

as u can see, the system will try to form more Ba(NO3)2 in order to 'remove' the excess nitrate. in doing so it uses up Ba2+.

but what happens to the BaSO4? its the precipitate, and it still is insoluble right? how does it redissolve?

Ba2+ (aq) + SO42- (aq) --> BaSO4 (s)

this is where the loss of Ba2+ comes in. in the first equation, we lost Ba2+ and nitrate to form barium nitrate. now the Ba2+ is in a shortage. by bla bla bla's principle, the equation will shift to the left = to produce more Ba2+ . and thus you witness the redissolution.

at least, that's what i think...  :P
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Offline breanainneire

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Re: precipitates re-dissolve in nitric acid
« Reply #2 on: January 13, 2007, 11:26:52 PM »
I guess that kind of makes sense. I've learned about the common ion effect before, but I don't really get why the Ba2+ would still be associating with the NO3-. Isn't the molecular formula when you add Ba(NO3)2(aq) and Na2SO4(aq):

Ba(NO3)2(aq) + Na2SO4(aq) --> NaNO3(aq) + BaSO4(s)  ?

Offline Donaldson Tan

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Re: precipitates re-dissolve in nitric acid
« Reply #3 on: January 13, 2007, 11:48:07 PM »
Ba2+ (aq) + 2NO3- (aq) --> Ba(NO3)2 (aq)

Ba2+ (aq) + 2NO3- (aq) = Ba(NO3)2 (aq)

I hate to spoil the fun here but the only re-dissolution of the precipitate upon addition of HNO3 should only happen to BaCO3 and Ag2CO3. This is purely an acid-base reaction.
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Offline xiankai

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Re: precipitates re-dissolve in nitric acid
« Reply #4 on: January 14, 2007, 12:03:30 AM »
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I don't really get why the Ba2+ would still be associating with the NO3-.

the reaction does not progress to completion, and hence there is tiny bit remaining.

Quote
Ba2+ (aq) + 2NO3- (aq) Ba(NO3)2 (aq)

i was trying to illustrate the equilibrium of the two ions among the others but i had no idea to go about it i guess...

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This is purely an acid-base reaction.

perhaps u can explain further? ;)
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Offline Borek

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Re: precipitates re-dissolve in nitric acid
« Reply #5 on: January 14, 2007, 05:53:41 AM »
Every precipitate is in equilibrium with some small amount of dissolved salt. Think what happens to CO32- anions when you add strong acid to the solution containing precipitate. What do you know about carbonic acid?

Note: IMHO Ag2CO3 should dissolve as well for exactly the same reasons.
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Offline breanainneire

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Re: precipitates re-dissolve in nitric acid
« Reply #6 on: January 14, 2007, 01:04:10 PM »
Sorry for my ignorance (i'm just in grade 12), but what should I know about carbonic acid?

Offline Borek

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Re: precipitates re-dissolve in nitric acid
« Reply #7 on: January 14, 2007, 01:25:11 PM »
I am more then sure that you will find the necessary information in your textbook.

What happens when you add strong acid to solid sodium bicarbonate? Or carbonate?
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Offline vhpk

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Re: precipitates re-dissolve in nitric acid
« Reply #8 on: January 14, 2007, 10:09:30 PM »
Yeah, of course, the strong acid will react with them and there is CO2, because, H2CO3 is a very weak acid, it'll be decomposited into CO2 and HO.
Another question: do you think that the precipitation of ion sulfur(like PbS or CuS) won't soluble in any acids
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Offline Borek

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Re: precipitates re-dissolve in nitric acid
« Reply #9 on: January 15, 2007, 03:38:58 AM »
because, H2CO3 is a very weak acid, it'll be decomposited into CO2 and HO.

It will decomposoe not because it is weak acid, but because it is unstable.

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Another question: do you think that the precipitation of ion sulfur(like PbS or CuS) won't soluble in any acids

Please stop adding new questions to old threads.
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Offline AWK

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Re: precipitates re-dissolve in nitric acid
« Reply #10 on: January 15, 2007, 04:51:47 AM »
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It will decomposoe not because it is weak acid, but because it is unstable.

Stability of H2CO3 depends on conditions.

Dry ether solution of carbonic acid can be evaporated at normal pressure, then carbonic acid sublimes, solidifies and as solid is stable at room temperature at prolonged time, but without a traces of water. water catalytically decomposes this acid.
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Offline Borek

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Re: precipitates re-dissolve in nitric acid
« Reply #11 on: January 15, 2007, 05:22:52 AM »
Yep. It can be also prepared by freezing water with carbon dioxide, radiating the mixture and then letting CO2 and H2O to sublimate. But for most practical purposes it is unstable.

Hi kids, please remember, that on the HS level carbonic acid is unstable always ;)
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