[galpinj]

I've read about dipoles and how the vectors cancel but haven't been satisfied with the reasoning. Supposedly, a molecule like CO2 (O=C=O) with two negative dipoles of equal strength will have its  dipoles cancel, making the molecule  nonpolar. However, I'm certain that each of the two oxygen atoms still retains its negative dipole and the carbon a positive dipole. So, Regardless of the vectors, these molecules remain polarized, and yet we call it nonpolar? Wouldn't each oxygen be attracted to positive ions? Isn't the electron densities uneven throughout the milolecule?

I've read that the charge remaining on the oxygens are defined as a quadrupole, but the answer is still somewhat unclear to me.

[mjc123]

Let's just clear up some inaccuracies. There's no such thing as a "negative dipole". A dipole is the separation of positive and negative charge. CO₂ has two dipoles equal in magnitude and opposite in direction, so the molecule overall has no dipole moment. But you are right that the charge distribution is non-uniform, although the symmetry of the distribution means there is no molecular dipole. The O atoms have a negative partial charge (not "dipole") and the C has a positive partial charge (not "dipole"). Each C=O bond is polar, but as you say, in the molecule as a whole the two dipoles cancel out. At distance, as in a gas, the molecule appears non-polar (there are quadrupole-quadrupole interactions, but they are much weaker than dipole-dipole), but at close quarters the local polarity of portions of the molecule becomes significant. Presumably that's how water reacts with dissolved CO₂ to give H₂CO₃ - the O of the water is attracted to the C of CO₂.

[galpinj]

Let's just clear up some inaccuracies. There's no such thing as a "negative dipole". A dipole is the separation of positive and negative charge. CO₂ has two dipoles equal in magnitude and opposite in direction, so the molecule overall has no dipole moment. But you are right that the charge distribution is non-uniform, although the symmetry of the distribution means there is no molecular dipole. The O atoms have a negative partial charge (not "dipole") and the C has a positive partial charge (not "dipole"). Each C=O bond is polar, but as you say, in the molecule as a whole the two dipoles cancel out. At distance, as in a gas, the molecule appears non-polar (there are quadrupole-quadrupole interactions, but they are much weaker than dipole-dipole), but at close quarters the local polarity of portions of the molecule becomes significant. Presumably that's how water reacts with dissolved CO₂ to give H₂CO₃ - the O of the water is attracted to the C of CO₂.

That really helps clarify some things, thank you! So, given what you have said, what is the significance of the term "dipole"? In our CO₂ example, the molecule is "nonpolar"; however, it still retains partial negative charges on the oxygen and a partial positive charge on the carbon (allowing it to interact with H₂O). If CO₂ can continue to interact using such intermolecular forces, then, to me, it seems like the molecule is still behaving as if it were polar! I would bet that these localized charges would even respond to electrical currents and other polar-related phenomena!

[Borek]

Matter of scale. For a molecule like CO₂, with two canceling dipoles, if you are far enough, the molecule looks like a neutral one, if you are close enough, there are two dipoles. In reality you are rarely close enough and small enough to see the CO₂ molecule as a polar one.

[Enthalpy]

The resulting quadripolar moment fades quickly over the distance, more so than a dipolar moment.

The partial charges have observable consequences: the vibrations of the gaseous CO₂ molecules absorb and emit IR strongly when they let the partial charges wobble with a net resulting electric current, for instance the V-shaped deformation.

On the other hand, the partial charges seem to contribute very little to the van de Waals' forces, even in contact, since CO₂ is about as soluble in polar and nonpolar solvents.

[mjc123]

If CO2 can continue to interact using such intermolecular forces, then, to me, it seems like the molecule is still behaving as if it were polar!

Polar but not dipolar. There is some separation of charge, but because of the symmetry the centre of the positive charge distribution coincides with the centre of the negative charge distribution, so there is no dipole moment. To illustrate, consider the distributions in the graph below. Both have a mean at x=0, but the blue one has a lot of its density near 0, the red one most of its density further away. If they were the distributions of positive and negative charge (I'm not saying they are in CO₂), you can see that there is some separation of charge (in fact two opposite dipoles cancelling out), but the centres of positive and negative charge coincide.

I would bet that these localized charges would even respond to electrical currents and other polar-related phenomena!

Depends what phenomena. For example, a dipole would tend to rotate in a uniform electric field, to line up with the field direction. A quadrupole doesn't. But a quadrupole can interact with a non-uniform electric field.
You need to get away from the dichotomy of "polar/non-polar", or thinking "polar" = "dipolar". It's more complicated than that.