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Dipole-dipole forces occur when the positive part of a polar molecule is attracted to the negative part of a polar molecule. In a nonpolar molecule, there may still be polar bonds, it's just that the dipoles cancel each other out. So why can't there be dipole-dipole forces between nonpolar molecules with polar bonds? There are still positive and negative parts of the molecule, so there can be attractions between them.

For example, in $\ce{CO2}$, there are negative areas near the oxygens and positive areas near the carbon.

electrostatic potential map of CO2

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Short answer: there are many electrostatic interactions between two non-polar molecules.

Beyond monopole (full charges) and permanent dipole moments (polar molecules), there is a full multipole expansion for the electrostatic potential around any molecule. (This is technically true for atoms and ions too, but higher-order terms are really only useful for molecules.)

So there are electrostatic potential energy interaction terms for charge-dipole, dipole-dipole, dipole-quadrupole, quadrupole-quadrupole, etc.

These terms are important - the quadrupole-quadrupole interactions dictate the orientation of the benzene dimer and $\ce{CO2}$ dimer in your example.1

The problem is that most of these interactions die off very quickly. The quadrupole-quadrupole term is:1

$$E(r)=\frac{-\Theta_1\Theta_2}{4\pi\epsilon_0r^5}\times\Gamma(\theta_1,\theta_2,\phi)$$

So roughly $1/r^5$, compared to $1/r^3$ for dipole-dipole interactions, or $1/r^6$ for dispersion forces like induced-dipoles.

When such molecules are close, the quadrupole moments (and other multipole electrostatic terms) can dictate packing and distances, but are not as strong or as long-range as dipole-dipole or charge interactions.

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  • $\begingroup$ So nonpolar molecules with polar bonds would have higher melting and boiling points than nonpolar molecules with no polar bonds right? $\endgroup$ – carbenoid Dec 30 '15 at 16:22
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    $\begingroup$ @swenger There are a lot of things that go into melting points and boiling points (e.g., internal rotations, etc.) but if you had two hypothetically equivalent non-polar molecules and one had polar bonds (and thus quadrupole interactions) it should have higher forces and higher mp and bp, yes. $\endgroup$ – Geoff Hutchison Dec 30 '15 at 17:11
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In the case of non-polar molecules, dispersion forces or London forces are present between them. These forces are induced dipole - induced dipole interactions. As we know that in non-polar molecule, the whole molecule has zero dipole moment but bonds are polar.

enter image description here

When two non-polar molecules comes closer to each other. The negative part (electrons) of one molecule attract the positive part (nucleus) of another molecule. As a result, two dipoles are induced. Such dipoles are called induced dipoles and interaction is called induced dipole - induced dipole interactions.

Example - CO2 - CCl4 and He - Ne intractions

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    $\begingroup$ This doesn't actually answer the question about whether there are dipole-dipole forces in molecules like $\ce{CO2}$ and if not, why. $\endgroup$ – bon Dec 30 '15 at 13:17
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Dipole-dipole interactions are electrostatic interactions between the permanent dipoles of different molecules. These interactions align the molecules to increase the attraction.

In order to form a dipole-dipole attraction, there should be a dipole moment for the considered molecule.

For an example: Water ($\ce{H2O}$)

It has a large permanent electric dipole moment. It's positive and negative charges are not centered at the same point; it behaves like a few equal and opposite charges separated by a small distance.

The permanent dipole in water is caused by oxygen's tendency to draw electrons to itself (i.e. oxygen is more electronegative than hydrogen). The 10 electrons of a water molecule are found more regularly near the oxygen atom's nucleus, which contains 8 protons. As a result, oxygen has a slight negative charge (δ-). Because oxygen is so electronegative, the electrons are found less regularly around the nucleus of the hydrogen atoms, which each only have one proton. As a result, hydrogen has a slight positive charge (δ+)

If we considered a NON POLAR molecule with POLAR BONDS as you mentioned; such as $\ce{CCl4}$ or $\ce{CO2}$

Molecules often contain polar bonds because of electronegativity differences but have no overall dipole moment if they are symmetrical. In the molecule tetrachloromethane ($\ce{CCl4}$), the chlorine atoms are more electronegative than the carbon atoms, and the electrons are drawn toward the chlorine atoms, creating dipoles. However, these carbon-chlorine dipoles cancel each other out because the molecular is symmetrical, and $\ce{CCl4}$ has no overall dipole movement.

Though $\ce{CO2}$ have polar bonds,it does not have a dipole moment, so it can not form dipole-dipole interactions.

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