[jjkwest1]

Carbon monoxide (CO) has a small dipole moment (μ = 0.11 D). Its bond length is 1.13 Å. While formaldehyde (H2CO) has a dipole moment of 2.3 D, the bond length of C=O is 1.21 Å. Explain why carbon monoxide has such a small dipole moment. What is the direction of the dipole moment of carbon monoxide you would predict?

Would the dipole moment of CO be small because the dipoles cancel out each other? Also would the dipole moment be towards the Carbon since it has a negative formal charge even though oxygen is the more electronegative atom?

[Jorriss]

No, it's not because the polarities cancel out, that can't happen with only two atoms that have vastly different electronegativities.

And it is because there is a negative charge on carbon...sort of... but that's not really a sufficient answer.

[jjkwest1]

I'm kind of lost on this, could you explain why the dipole is small and if the net dipole moment points to the carbon atom is correct?
Thanks

[jjkwest1]

Also what i meant to say was that since the dipole moment points to both carbon and oxygen they make themselves smaller thus making the dipole moment small but i don't know if thats correct.

[Jorriss]

Maybe you had the right idea but what you wrote wasn't quite correct.

What it means that polarities cancel out is you have two atoms of equal, or near equal electronegativity, covalently bonded to eachother.

For example, the simplest case, H2. H-H, Both hydrogens share the same electrons but they 'want' the electrons equally so their electronegativities cancel out, making a nonpolar bond.

What it means when you say dipoles cancel out is like in the case of CO2.

O=C=O

Each C=O bond is polar in the direct of oxygen, but there are two equal bonds on opposite sides of each other so they completely cancel out. This is the case of dipoles canceling out.


Now lets look at carbon monoxide.

:C=O::


Anyways, more directly on the problem.

To explain the small dipole moment you'll have to incorporate two concepts - resonance and inductive effects.

[jjkwest1]

Thanks Jorriss, lastly would it be correct to assume that the net dipole is pointed towards the carbon?

[Jorriss]

Without being told explicitly I wouldn't know.

Look at the structure again

:C=O::

There is a resonance structure that gives carbon a negative formal charge and oxygen a positive formal charge. But we have to consider how good is the resonance structure. It creates separation of charge and adds formal charges for a complete octet on carbon. It's probably decent.

But at the same time, oxygen is more electronegative than carbon meaning electron density is going to be pulled in the direction of oxygen, away from carbon.

So now we have these two competing factors, the resonance which suggests a dipole in the direction of carbon, and inductive effects, which suggests a dipole facing oxygen.

Is the resonance worth more than induction in this case? If the dipole is .11 D, it seems the answer is they're actually comparable. I would wager the dipole is in the direction of oxygen myself, but the important thing to take away from this question is the reasoning behind the choice, not the choice itself.

[Jorriss]

I can't edit my post.

'What it means that polarities cancel out is you have two atoms of equal, or near equal electronegativity, covalently bonded to eachother.'

That really should read 'What it means for a bond to possess no dipole,' as the polarities aren't canceling out there.