# Why doesn't CO2 repel itself? [closed]

In a molecule like $\ce{CO2}$, the oxygens are more electronegative than the carbon, and so the electron cloud is denser around them. Then, isn't the oxygen going to be a negative dipole? The picture I have in my mind is a positive magnet sandwiched between two negative magnets. So my question is – why don't $\ce{CO2}$ molecules repel each other?

## closed as unclear what you're asking by Klaus-Dieter Warzecha, Jon Custer, NotEvans., Todd Minehardt, Geoff HutchisonDec 23 '16 at 3:07

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• Related: Why is O=C=O nonpolar? – Loong Dec 22 '16 at 1:56
• There is no such thing as negative dipole. As for the molecules repelling each other: why, they surely do, much like any other stable molecules. – Ivan Neretin Dec 22 '16 at 7:52

Carbon dioxide is a linear molecule. Therefore, you have two directly opposite dipoles (each point from the C to one O). These cancel each other out and CO2 is actually nonpolar, as detailed in the answer to this question. Then, two CO2 molecules can interact and attract each other through london dispersion forces.

That all said, even in the hypothetical case where those dipoles didn't cancel each other out, and you somehow did have electron density at two different ends of the molecule, I still wouldn't expect CO2 molecules to all repel each other. In that thought experiment, your carbon would still be strongly positively-charged, so you would still have multiple CO2 molecules clustering via the electron density on the oxygens interacting with the partial positive charge on the carbons of other molecules.

Carbon dioxide has no permanent electric dipole. Atoms do not behave like the small bar magnets you suggest. However, the answer your question 'why don't $\ce{CO2}$ molecules repel each other' is both 'yes' they do and 'no' they don't as I try to explain.

Between one atom and another or between two molecules there is always a dispersion force even when there is not a permanent dipole. Although quantum mechanical, the origin of this force can be understood intuitively as follows. Even though the average dipole moment is exactly zero, at any instant in time, the positions of the electron may not be exactly symmetrical and so an instantaneous dipole is produced. This dipole generates an electric field that propagates at the speed of light and polarises nearby atoms or molecules, and this induces a second instantaneous dipole. The resulting induced-dipole induced-dipole interaction produces an attractive force that has a finite time-average i.e. it is not zero.

Experiments on many gases, for example $\ce{CO2}$, show that the ideal gas law is not obeyed exactly. To describe the experimental data van-der-Waals extended the ideal gas law to become $(p-a)(V-b)=RT$ where b accounts for the finite volume of a real gas and a accounts for the attractive potential, one cause of which is described.

At a given temperature there is a distribution of molecule speeds, (or energies) as described by the Maxwell-Boltzmann distribution. A few molecules have a low speed, and a few very high speeds, and obviously, most fall between these limits. The distribution looks sort of like a lopsided bell shaped curve. When two molecules collide and they are travelling relatively slowly wrt one another then their mutual energy may not be sufficient to overcome the attractive dispersion energy and so may pair up for a short while until, perhaps, the pair collide with another molecule and separate. This reduces the overall pressure because it reduces the effective number of molecules in the gas that can collide with the walls.

When the relative collision energy is far larger than the attractive dispersion energy, the potential is no longer limiting and the molecules can exchange energy and momentum as collide one another and separate.