[offlinedoctor]

Hi,

I'm trying to understand strengths of acids by their stability when turning into anions, I'm just wondering what exactly do they mean by, the more stable the anion the stronger the acid. Does that mean, the less negative it is when deprotonated, the more likely it is to not accept hydrogen in its conjugate base, and hence, the stronger the acid?

Also, for the following reaction (in attachments), I have trouble seeing how the methyl amine donated its electrons if it has an octet in the final product, yet, no electron loss.

Thanks!

[Hunter2]

Boron has a electron gap. So the non bonding electrons of the nitrogen can jump in.

[orgopete]

Let me suggest an alternative model for acidity. In the H_nX acids, the bond lengths decrease from C to F, while the acidity increases. You could consider that as the electrons are pulled toward the nucleus, they are also pulled away from the proton. An increase in distance results in an increased acidity. Basicity is the opposite. The further the electrons extend from the nucleus, the easier they can abstract a proton. Hence, the stronger the acid (the more the electrons are pulled in), the weaker the conjugate base. (I think it is actually a little more complicated than that, but as a model to think of, it can work.)

For the complex, we don't subtract electrons in calculating whether an octet is present, though we do for formal charge. Formal charge is a bookkeeping device, charges don't change. Protons are positive, electrons negative. Nitrogen is +7 on the left and the right. However, we can think that the donated electrons of the nitrogen complete the octet of both boron and nitrogen. Would you be surprised to know BH4(-) has the longest bonds of the completed octets in its row? Why might BH4(-) donate this additional pair of electrons with a proton in some reactions?

[TwistedConf]

I'm trying to understand strengths of acids by their stability when turning into anions, I'm just wondering what exactly do they mean by, the more stable the anion the stronger the acid. Does that mean, the less negative it is when deprotonated, the more likely it is to not accept hydrogen in its conjugate base, and hence, the stronger the acid?

You've kind of got it.  Increasing conjugate base anion stability means the ion is more stable and less basic, and thus it would have a stronger and stronger conjugate acid.

What you need to work on is how to evaluate the stability of the conjugate bases by looking at them. There are several factors involved and your statement "the less negative it is" has issues.

[TwistedConf]

Why might BH4(-) donate this additional pair of electrons with a proton in some reactions?

What are you talking about?  There's no lone pair of electrons in BH₄^-.

[orgopete]

Why might BH4(-) donate this additional pair of electrons with a proton in some reactions?

What are you talking about?  There's no lone pair of electrons in BH₄^-.

Agreed, no lone pair. Borohydride is a hydride donor, sort of. It probably isn't a hydride that gets donated, it is the electrons between the boron and hydrogen. If you protonate the electrons, the electrons remain with the proton to give H2 and a trivalent boron.

I was suggesting something about the nature of the elements. Acidity and basicity is about how electrons are held. As they are pulled closer to a nucleus, the become weaker bases. As they extend further, the become stronger bases. There is a limit though as the atoms no longer hold the electrons and they become electron donors. If we are comparing hydrides, LiH, BeH2, and BH3 (via ROBH₃-) are hydride donors.

Ah, well, you are right. This is becoming too esoteric. It is simpler to just check a pKa table.