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Topic: Pka of Amines  (Read 64728 times)

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Offline mrlucky0

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Pka of Amines
« on: February 28, 2008, 11:43:02 PM »
Can someone explain to me why a methylamine, CH3NH2 (Pka = 10.64) is more basic then a trimethylamine, (CH3)3-N (Pka = 9.87)? From what I know, I thought less electronegative substituents, like alkyl groups increase the basicity. Hence, shouldn't trimethylamine be more basic because it's got 3 methyl groups?

Offline refid

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Re: Pka of Amines
« Reply #1 on: February 29, 2008, 01:15:49 PM »
I believe basicity is determined by the availability of the lone pare on the nitrogen.

Your right about the methyl substituent are less electronegative, but do draw in electrons inductively (-I effect) from the nitrogen.

So if you think about it:
1 methyl substituent draws let say 0.1% of electrons (methylamine),
3 methyl substituent must draw 0.1% x 3 = 0.3% of electrons (trimethylamine).

This should explain why trimethylamine is less basic then methylamine.

Hope this helps.

Offline Rico

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Re: Pka of Amines
« Reply #2 on: February 29, 2008, 05:43:01 PM »
Hey mrlucky0 and refid

I must say, refid, that i do not agree with you - methylgroups in amines, and ammonium-ions for that matter, are electrondonating and not electronwithdrawing.

The basicity trends of amines are most easily explained by looking at the acidity of ammonium-ions, which (in solution) follows the trend given below:

NH4+ > (CH3)3NH+ > CH3NH3+ > (CH3)2NH2+

That is NH4+-ions are the most acidic (least basic) ones, whereas (CH3)2NH2+-ions are the least acidic (most basic) ones.

This trend can be explained by two conflicting effects - polarizability and solvation.

If you look at gas phase acidities the trend is as follows:

NH4+ > CH3NH3+ >  (CH3)2NH2+  > (CH3)3NH+

Which is what we would normally expect since more methyl-groups donates more electrons spreading out the positive charge and therefore stabilizing the ammonium-ion (making it less acidic, and the conjugate base more basic).
On the other hand solvation effects leads to the following trend:

(CH3)3NH+ > (CH3)2NH2+ > CH3NH3+ > NH4+

This trend can be explained by the size of the ammonium-ion - the larger, the least well solvated, therefore least stable and hence more acidic (less basic).
A combination of these two trends gives the acidity (and therefore basicity) trend of ammonium-ions (amines) in solution.

I hope this answers your question!

Rico

Offline movies

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Re: Pka of Amines
« Reply #3 on: February 29, 2008, 06:48:51 PM »
Excellent explanation, Rico!  I agree completely!

Offline AWK

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Re: Pka of Amines
« Reply #4 on: March 03, 2008, 06:36:39 AM »
Can someone explain to me why a methylamine, CH3NH2 (Pka = 10.64) is more basic then a trimethylamine, (CH3)3-N (Pka = 9.87)? From what I know, I thought less electronegative substituents, like alkyl groups increase the basicity. Hence, shouldn't trimethylamine be more basic because it's got 3 methyl groups?
Trimethylamine is more basic then methylamine. Your data are pKas that say directly about acidities of ammonium cations
AWK

Offline refid

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Re: Pka of Amines
« Reply #5 on: March 05, 2008, 12:27:58 AM »
Opps..Thanks rico

Offline adam

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Re: Pka of Amines
« Reply #6 on: March 05, 2008, 04:21:21 AM »
yes AWK, the pka values are the pka's of their conjugated acids,
-the highest pka the stronger base? it is? methylamine should be more basic than trimethyl amine.

the pka's value of alkyl amines are around 30????

thanks!

Offline AWK

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Re: Pka of Amines
« Reply #7 on: March 05, 2008, 09:06:48 AM »
AWK

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