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Why is the solution of ammonium bifluoride more acidic than aqueous ammonium fluoride?
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Magnus Wootton
Why is the solution of ammonium bifluoride more acidic than aqueous ammonium fluoride?
This question and the one referenced by Mithoron address $\ce{NH4HF2}$ as an entity in itself which must be examined as a whole.
It is easier conceptually to rewrite the formula for the compound as $\ce{NH4F·HF}$. Consider adding $\ce{NH4F}$ to water: you get a $\mathrm{pH}$ near $7.$ In a separate container, add $\ce{HF}$ to water (approx. $\pu{0.1 M});$you get a $\mathrm{pH}\approx 1$.
Combine the solutions and the $\mathrm{pH}$ will settle out somewhere between $1$ and $7,$ because of the common ion effect: $\ce{F-}$ from the $\ce{NH4F}$ inhibits the ionization of $\ce{HF},$ but the solution is still acidic. The approximate $\mathrm{pH}$ can be calculated from the ionization constant of $\ce{HF}$ and the concentrations of $\ce{NH4F}$ and $\ce{HF}.$
The concept of the $\ce{HF2-}$ moiety dominating the discussion for a solution is probably making the analysis murkier than it needs to be. It's just a hydrogen-bonded complex that happens to retain some semblance of existence in the solid. It is fairly stable: mp $\pu{126 °C},$ bp $\pu{240 °C}$ (with decomposition).
This question and the one referenced by Mithoron address $\ce{NH4HF2}$ as an entity in itself which must be examined as a whole.
It is easier conceptually to rewrite the formula for the compound as $\ce{NH4F·HF}$. Consider adding $\ce{NH4F}$ to water: you get a $\mathrm{pH}$ near $7.$ In a separate container, add $\ce{HF}$ to water (approx. $\pu{0.1 M});$you get a $\mathrm{pH}\approx 1$.
Combine the solutions and the $\mathrm{pH}$ will settle out somewhere between $1$ and $7,$ because of the common ion effect: $\ce{F-}$ from the $\ce{NH4F}$ inhibits the ionization of $\ce{HF},$ but the solution is still acidic. The approximate $\mathrm{pH}$ can be calculated from the ionization constant of $\ce{HF}$ and the concentrations of $\ce{NH4F}$ and $\ce{HF}.$
The concept of the $\ce{HF2-}$ moiety dominating the discussion for a solution is probably making the analysis murkier than it needs to be. It's just a hydrogen-bonded complex that happens to retain some semblance of existence in the solid. It is fairly stable: mp $\pu{126 °C},$ bp $\pu{240 °C}$ (with decomposition).
@andselisk: Well, thanks, I guess. Its not clear to me what straightforward notation, according to my screen, gets retranslated to a different format. I have already been notified about certain subscripts, but not the temperature. Is there a way that I can see the wrong view before I post it?More
"126$^0$ C" reads as "one hundred twenty-six to the power of zero coulombs". "NH$_4$F.HF" has a dot that doesnt make sense and the formula breaks across lines incorrectly. Please consider using proper formatting tools (More
on how to format your future posts better with MathJax and Markdown. There are no other tools other than the preview windows, standardized notations (which everyone is supposed to look up) and common sense, Im afraid.More
This question and the one referenced by Mithoron address $\ce{NH4HF2}$ as an entity in itself which must be examined as a whole.
It is easier conceptually to rewrite the formula for the compound as $\ce{NH4F·HF}$. Consider adding $\ce{NH4F}$ to water: you get a $\mathrm{pH}$ near $7.$ In a separate container, add $\ce{HF}$ to water (approx. $\pu{0.1 M});$ you get a $\mathrm{pH}\approx 1$.
Combine the solutions and the $\mathrm{pH}$ will settle out somewhere between $1$ and $7,$ because of the common ion effect: $\ce{F-}$ from the $\ce{NH4F}$ inhibits the ionization of $\ce{HF},$ but the solution is still acidic. The approximate $\mathrm{pH}$ can be calculated from the ionization constant of $\ce{HF}$ and the concentrations of $\ce{NH4F}$ and $\ce{HF}.$
The concept of the $\ce{HF2-}$ moiety dominating the discussion for a solution is probably making the analysis murkier than it needs to be. It's just a hydrogen-bonded complex that happens to retain some semblance of existence in the solid. It is fairly stable: mp $\pu{126 °C},$ bp $\pu{240 °C}$ (with decomposition).
This question and the one referenced by Mithoron address $\ce{NH4HF2}$ as an entity in itself which must be examined as a whole.
It is easier conceptually to rewrite the formula for the compound as $\ce{NH4F·HF}$. Consider adding $\ce{NH4F}$ to water: you get a $\mathrm{pH}$ near $7.$ In a separate container, add $\ce{HF}$ to water (approx. $\pu{0.1 M});$ you get a $\mathrm{pH}\approx 1$.
Combine the solutions and the $\mathrm{pH}$ will settle out somewhere between $1$ and $7,$ because of the common ion effect: $\ce{F-}$ from the $\ce{NH4F}$ inhibits the ionization of $\ce{HF},$ but the solution is still acidic. The approximate $\mathrm{pH}$ can be calculated from the ionization constant of $\ce{HF}$ and the concentrations of $\ce{NH4F}$ and $\ce{HF}.$
The concept of the $\ce{HF2-}$ moiety dominating the discussion for a solution is probably making the analysis murkier than it needs to be. It's just a hydrogen-bonded complex that happens to retain some semblance of existence in the solid. It is fairly stable: mp $\pu{126 °C},$ bp $\pu{240 °C}$ (with decomposition).
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