Club Chemistry: The Anomeric Effect

In a post several months ago, I was talking about sugars and mentioned:

Note that in the cyclic isomer of glucose – β-D-glucopyranose (left) – all 5 substituents on the pyran ring are in the low-energy equatorial position (actually, the lowest-energy conformation of glucose is α-D-glucopyranose, where one of the -OH substituents is in the axial position. It is stabilized by what is known as the anomeric effect)

To which Mitch commented:

In regards to the anomeric effect, no one else finds it strange that when there is no good stereoelectronic effect explanation all of a sudden hyperconjugation is the key?

First, a mea culpa. For unsubstituted glucopyranose, the β isomer is the lowest-energy isomer, and the α isomer is disfavored at a ratio of about 64:36. When the hydroxyl group at the acetal position is changed to a methoxy group, then the α isomer is the lowest-energy isomer at a ratio of about 67:33 – the selectivity reverses. (click images for larger throughout) (update: figure labels fixed)

Second, my PhD research relies heavily on the anomeric effect, and I often get this ‘I don’t understand the anomeric effect’ response from people. They assume it’s all handwaving. I’d like to explain the anomeric effect and hopefully clear up some of the confusion surrounding it. Read more below the jump.

We remember from undergrad organic classes, that substituents on a cyclohexane ring prefer the equatorial position to relieve steric strain. There is a corresponding ring-flipped chair in which the substituent is axial, and the energy barrier associated with this ring flip can be quantified. Forcyclohexanol, the equatorial position is favored at a ratio of about 9:1.

When an endocyclic heteroatom (typically oxygen, nitrogen, or sulfur) is introduced at the position adjacent to the hydroxy group, for the example of methoxypyran, the hydroxy substituent now prefers the axial orientation – in spite of the associated steric strain – at a ratio of about 4:1. This trend is not limited to oxygen substituents. Any electronegative element will prefer an axial position when a heteroatom is on the adjacent endocylic position. For the fluoro-xylose derivative in the image, the fluorine prefers the axial position EVEN THOUGH all of the other substituents are also axial.

This phenomenon is known for acylic systems, too. If you were asked to draw a Newman projectionfor pentane and dimethoxymethane, for pentane you would drawn the Newman projection where the two bulky substituents are antiperiplanar to each other to minimize steric interactions. And you’d be right. But for dimethoxymethane, the two bulky substituents are gauche to each other in the lowest-energy Newman projection. As a hint to what’s going on, this puts a lone pair on oxygen antiperiplanar to the electronegative substituent.

So why is this happening? There are two main explanations given, and they both work together to explain why the seemingly more sterically hindered conformation is the most stable. On one hand, when the electronegative substituent is in an equatorial orientation, the local dipole moment of the substituent and the local dipole of the endocyclic heteroatom are pointing in relatively the same direction. This alignment of dipoles leads to a large net dipole for the molecule. However, when the electronegative substituent is axial, the local dipoles are more or less pointing away from each other. This relief of a net dipole is stabilizing for the molecule. This is a fine explanation, but I don’t think it fully accounts for all of the stabilization seen in these systems.

The other main explanation is a molecular orbital argument. In short, whenever you can lower the energy of a system, the system is net more stable. When the electronegative substituent is in an axial orientation, the C-X sigma* antibonding orbital is directly lined up with the axial lone pair of electrons on oxygen. This allows for some delocalization of the lone pair of electrons into the sigma* antibonding orbital. The electrons from the lone pair interact with the sigma* antibonding orbital in a stabilizing manner. The elctron density is now spread over two atoms. The delocalization of the electrons results in a stabilization of those electrons, and leads to a net stabilization of the molecule.

But why don’t all groups next to an endocyclic heteroatom prefer the axial orientation, you might ask? Why wouldn’t a regular methyl group be stabilized in the same way? There is a trend amond endocyclic and exocylic groups which rates the ability to donate or accept electrons. A carbanion is among the best electron donors, and heteroatom lone pairs are pretty good donors, too. Things like C-H sigma bonding electrons aren’t such good donors, but do delocalize to some extent (this is why teritary carbocations are more stable than methyl carbocations). There is a similar trend for electron accepting groups. An empty p orbital is among the best acceptors, and C-X sigma* antibonding orbitals are pretty good, too. The reason for this is the relative energy of the donating and accepting orbitals. Carbanions and heteroatom lone pairs are relatively higher in energy than C-H sigma bonding electrons, and p orbitals and C-X sigma* antibonding orbitals are relatively lower in energy than C-C sigma* antibonding orbtials. When the donor and acceptor orbitals are closer in energy, the stabilization is more favorable.

I talked about this general phenomenon as it relates to hyperconjugation over on the forums, too, if you want to read more.

So to summarize, for a cyclic system, when an endocyclic atom (Y) has a lone pair of electrons, a neighboring electronegative substituent (X) prefers to reside in an axial orientation. More generally for cyclic or acylic systems, when a heteroatom (Y) has a lone pair of electrons, an neighboring electronegative group prefers a gauche orientation – in spite of what sterics might dictate. This allows for maximum overlap of the lone pair of electrons and the neighboring C-X sigma* antibonding orbital.