[cseil]

Hello everyone,
I have problems comparing the basicity of these compounds.



dimethyl ether



dimethyl sulfide



cyclohexylamine



N-methylbutylamine



aniline

I have to write them in order of basicity.

First of all, the most electronegative bond is C-O, then there's C-N and C-S.
So I believe that dimethyl ether is the most basic compound in here.
It has two methyls (electron donating group) and an electronegative bond. So the electrons on oxygen are more available.

Then there are the amines.
N-methylbutylamine is a secondary amine. It is more basic than cyclohexylamine (primary).
I can also say that N-methylbutylamine has a carbon with two electron donating groups, while cyclohexylamine has only one.

After amines, I believe there's dimethyl sulfide and then aniline.
Aniline is stabilised by resonance and it is the most acidic one, so the least basic.

Is that right?
Where do I get wrong, if I do?

[Babcock_Hall]

Your statement about aniline does not make sense to me.  How do you know it is most acidic, for example?  Also, your assumption about the relationship between basicity and electronegativity is certainly open to question.  Electronegativity and acidity are poorly correlated in a number of instances.

[cseil]

If I consider the conjugated base of aniline,
it is



It is stable because of the resonance. So it is a weak base with a relatively strong conjugated acid.
That's why I said that aniline is the most acidic.

I don't know if it is right.

Talking about the relationship between basicity and electronegativity, I think I should consider it the opposite way. Electronegative means more acidic according to my book, so the speech about C-O, C-N and C-S is not like I said in the first place, but the opposite.

Why do you say that electronegativity and acidity are poorly correlated? There are pages on my book that says that.

[orgopete]

.

Talking about the relationship between basicity and electronegativity, I think I should consider it the opposite way. Electronegative means more acidic according to my book, so the speech about C-O, C-N and C-S is not like I said in the first place, but the opposite.

Why do you say that electronegativity and acidity are poorly correlated? There are pages on my book that says that.

A reason to say acidity and electronegativity are poorly correlated is if you compare CH4, NH3, H2O, and HF, HF is the most electronegative and the most acidic. If you compared HF, HCl,  HBr, and HI, HF is again the most electronegative, but the least acidic.

The basicity order compares the pKa of the conjugate acids. How easily can a sulfur be protonated compared to an oxygen or nitrogen. H3S+ is more difficult to find, but the acidity of H2S and H2O will give the order. For the other compounds, I suggest looking up the pKa of the conjugate acids. Then, as you have been suggesting their order, you can check your explanations. If this requires you to adjust how you think about electron availability, it will prove useful here and elsewhere.

[cseil]

I knew HI is the most acidic between HF, HCl, HBr because of the dimension and the polarizability.
In that case electronegativity does not count.


Electronegativity does not count in the comparison between dimethyl ether and dimethyl sulfide.
In this case, for the same reason I would say that H2S is more acidic than H2O, so dimethyl sulfide is less basic than dimethyl ether.


pKa of cyclohexylamine is 10.45
pKa of N-methylbutylamine is 10.55.

N-methylbutylamine is more basic than cyclohexylamine. I said why: it has two electron donating groups, one is a secondary amine, the other is a primary amine.

My problem now is how to compare the basicity of oxygen/sulfur compounds with the nitrogen compounds.
If I think about electronegativity, the C-O bond is more electronegative, that means more acidic. So the dimethyl ether is less basic than the two compounds with nitrogen.

The scale could be: N-methylbutylamine > cyclohexylamine > dimethyl ether > dimethyl sulfide > aniline

Last part I don't know if that's right

[Babcock_Hall]

If you knew how acidic Ph-NH₃⁺ were, then you could say something meaningful about how basic its conjugate base (aniline) is.

[cseil]

anilinium ion has a pka 4.6.
it is lower than H2O, H2S, NH3.

[Babcock_Hall]

The pK_a of H₃O⁺ should be compared with the pK_a of the anilinium ion, Ph-NH₃⁺.  They are similar or identical to the conjugate acids of the bases you are comparing.

[cseil]

So here's my mistake.
I have considered well the ether, sulfide and amines.

The problem is aniline. pKa of H3O+ is -1.74. I expect H3S+ pka to be even lower and the pKa of NH4+ to be higher of both (NH4+ is 9.24).
pKa of anilinium ion 4.6. So it is less acidic than H3O+ and H3S+ but more acidic than NH4+.

So the scale of my compounds is:

N-methylbutylamine > cyclohexylamine > aniline > dimethyl ether > dimethyl sulfide

Am I right?
Thank you, things are gettin clearer to me now.
I always considered aniline stabilised by resonance, and talking about acidity I was right (because I considered the conjugated base). But now I'm considering the basicity, so I have to consider not the conjugated base, but the conjugated acid. And it is a little bit different. Ph-NH3+ doesn't have any resonance effect.

[Babcock_Hall]

The table I am using gives pK_a of -2 to -5 for protonated alcohols (R-OH₂⁺ and -7 for a protonated alkyl thiol).  Therefore, the protonated thiol is the stronger acid.

Aniline is an interesting case.  How do you think that resonance affects the free base form?

[cseil]

What do you mean with "free base form"? I'm sorry but I don't know. 

[Babcock_Hall]

The neutral form (aniline) is the free base.  The conjugate acid is the anilinium ion.  It must have a counterion and therefore is not free.  The pK_a of the anilinium ion is roughly four pH units removed from the conjugate acids of alkyl amines (sorry I don't have the pK_a in front of me at the moment).  The question is why it is a much stronger acid, or conversely why aniline is a much weaker base.

[cseil]

The neutral form (aniline) is the free base.  The conjugate acid is the anilinium ion.  It must have a counterion and therefore is not free.  The pK_a of the anilinium ion is roughly four pH units removed from the conjugate acids of alkyl amines (sorry I don't have the pK_a in front of me at the moment).  The question is why it is a much stronger acid, or conversely why aniline is a much weaker base.

I think it is because the anilinium ion has a positive charge and nitrogen can't stabilise it in any way. This instability reflects the reactivity of the Ph-NH3+. It reacts easily because the conjugated base is aniline, a very stable compound compared with anilinium ion. So, the instability of anilinium makes it a stronger acid, the stability of aniline makes it a weak base.

I can't really figure out why Ph-NH3+ is so much a stronger acid than alkyl amines.
I suppose there's some kind of influence from the electrons of the aromatic ring.

I haven't studied the chapter about aromacity yet   

[Babcock_Hall]

Are there more than one resonance structures you can draw for aniline?

[cseil]

Are there more than one resonance structures you can draw for aniline?

Yes, there are three of them.
So here's the big difference with the alkyl amines?