Chemical Forums
Chemistry Forums for Students => Organic Chemistry Forum => Topic started by: minimal on September 15, 2008, 04:04:38 PM
-
Hi, I was told that amine solubility in water is hindered at high pH levels. I do not understand this, as deprotonating the amine would still leave it available for hydrogen bonding.
Can someone shed some more light (and resources) on this?
-
What is more polar - amine, or protonated amine (cation)?
-
regular amine? Wouldn't it depend on whether or not the ammonium ion were free or not? As in, any ionic bonds it would form would obviously be polar, but by itself it just represents a +1 charge, correct?
-
What is being protonated?
-
What is being protonated?
in your question or the original? In original I'm talking about de-protonation.
However, the nitrogen of the amine is being protonated in your question, removing the bent tetrahedral conformation by attacking the lone pair and making it closer to the standard 109 degree angle.
-
However, the nitrogen of the amine is being protonated in your question, removing the bent tetrahedral conformation by attacking the lone pair and making it closer to the standard 109 degree angle.
As you have correctly wrote - lone pair. Each nitrogen in each amine has a lone pair, so every amine is protonated in exactly the same way, and in each case you end with just a cation - which, as an ion, is much better soluble. Now, how easily it gets protonated (how strong base it is, or what pKa protonated amine has) depends on the many factors, but overall effect is always similar - in low pH amine gets protonated and becomes better soluble, in high pH it gets deprotonated and the effect on solubility is opposite.
-
I still don't understand why that means the amine would become insoluble. I can understand it becoming slightly less soluble. But if you were trying to recover any of the product by adding an organic solvent, very little would migrate afaik.
-
I'm trying to work out here why there's less solubility in unprotonated or deprotonated amines. Do you think it's because the hydrogen of water want to hydrogen bond with nitrogen, but when it has to do it against a lone pair there's more steric hindrance factors at work preventing it from doing so? However, when there's a substantial bond length (as is the case in ammonium or other positively charged amines), it is able to form the hydrogen bond easier?
-
I think you are overcomplicating. Charged molecule is more polar, thus it is more likely to dissolve being surrounded by water dipoles (solvated if you like).
-
I think you are overcomplicating. Charged molecule is more polar, thus it is more likely to dissolve being surrounded by water dipoles (solvated if you like).
but by that logic, an amine that has been severely depronated the NH2- anion, should be very soluble, no?
-
Yes.
Although I suppose in most cases it will immediately react with water. But I can be wrong. Any organiker around?
-
ok, one more question, why is the charged molecule more soluble than a simple polar molecule? Is it able to hold the partially negative charged aqueous oxygens in close proximity easier, while also likely repelling other NH4+ species away?
Is this because now each hydrogen of the amine still possess a partial positive, while the nitrogen also has a positive formal charge?
Or do the hydrogens lose their partial positives when the nitrogen possesses a formal charge?
-
I suppose in most cases it will immediately react with water
Yes, it will immediately react with water to form the neutral amine.
Looking back through the thread, I think you two may have been talking about different things. There are three possibilities for the nitrogen atom of an amine at varying "pH":
LOW pH: The amine will be protonated - NR3H+. This ammonium cation will be water soluble.
MID pH: Careful with your wording here. The amine will be the neutral NR3. The amine (NR3) will be 'deprotonated' with respect to the ammonium cation (NR3H+), but the amine (NR3) will be 'protonated' compared to its anion (NR2-).
HIGH pH: The amine will be deprotonated and will exist briefly as its metal amide (NR2-). As noted, though, if we're talking about an aqueous solution here, the metal amide will be the strongest base and will pull a proton off water to protonate the metal amide (NR2-) to the neutral amine (NR3).
I think when minimal was using the word 'deprotonated' (as in the original post), the term was being used to mean the product of the reaction would be NR2-. When Borek was using 'deprotonated', it was used to mean the product of the reaction would be NR3. Am I interpreting your intentions wrong?
Bottom line: NR3H+ is water soluble. NR3 is not water soluble (edit: this is inaccurate, see clarification in subsequent post). NR2- will react with water to form NR3 - not water soluble.
-
I should amend my remarks, as they're not exactly accurate.
Simple amines are water soluble. Ammonia is water soluble (it's the basis for simple glass cleaners - solutions of ammonium hydroxide). Think of them almost like alcohols - simple alcohols are water soluble (mainly due to hydrogen bonding as you noted earlier), but as the alkyl chain becomes longer, the molecule becomes more hydrocarbon like and less alcohol like, and the solubility diminished as a result until alcohols with long alkyl chains are no longer water soluble.
Same thing with amines. As the length (and number) of alkyl groups grows, the amine becomes less and less water soluble until it is no longer water soluble.
Amines that are basic enough will - to some extent - deprotonate water to form the ammonium cation, and the ammonium cation will be water soluble. Similarly, all hydrochloride salts of amines (the corresponding ammonium chloride) will dissolve in water.
Hope this clears things up. Sorry my first post was a tad inaccurate.
-
I should amend my remarks, as they're not exactly accurate.
Simple amines are water soluble. Ammonia is water soluble (it's the basis for simple glass cleaners - solutions of ammonium hydroxide). Think of them almost like alcohols - simple alcohols are water soluble (mainly due to hydrogen bonding as you noted earlier), but as the alkyl chain becomes longer, the molecule becomes more hydrocarbon like and less alcohol like, and the solubility diminished as a result until alcohols with long alkyl chains are no longer water soluble.
Same thing with amines. As the length (and number) of alkyl groups grows, the amine becomes less and less water soluble until it is no longer water soluble.
Amines that are basic enough will - to some extent - deprotonate water to form the ammonium cation, and the ammonium cation will be water soluble. Similarly, all hydrochloride salts of amines (the corresponding ammonium chloride) will dissolve in water.
Hope this clears things up. Sorry my first post was a tad inaccurate.
do you know where I might find an equation relating london forces to hydrogen bonding, so that I might be able to predict at what alkyl chain point the amine becomes soluble in either non-polar/polar liquids?
-
Not a quantitative list, per se, but the following sites seem to indicate the water solubility trends of primary amines roughly follow alcohols, and tertiary amines roughly follow ethers, with secondary amines somewhere in the middle. At least two of the sites specifically mentioned that water solubility drops significantly after alkyl chains of 6 carbon atoms or more.
http://en.wikipedia.org/wiki/Amino
http://members.aol.com/logan20/amines.html
http://www.chemguide.co.uk/organicprops/amines/background.html
http://dl.clackamas.cc.or.us/ch106-05/properti.htm
http://www.cem.msu.edu/~reusch/VirtualText/amine1.htm
-
Great, thanks for the help both of you.
I think the problem was I didn't realize that @ pH of 7, NH4 exists instead of NH3. The whole problem came about because I read someone elses suggestion of using NaOH and then organic solvent to remove a polyamine. I assumed that polyamine would have already been NH3 (or equivalent obviously NR3), and that adding NaOH would deprotonate it to NH2-. But I realize now that that would not exist.
I have a followup question in light of all this...I was always told that a strong acid has a weak conjugate base, and vise versa. Why is it then that NH4 is a weak acid, but NH3 is also a weak base? Is it just a very loose rule that a strong acid has a weak conjugate base, and a weak acid has a strong conjugate base?
-
Is it just a very loose rule that a strong acid has a weak conjugate base, and a weak acid has a strong conjugate base?
This is a very precise rule:
pKa + pKb = 14
Trick is, terms "weak acid" and/or "weak base" are loosely defined.