Can nonpolar molecules exhibit dipole-dipole forces?

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](https://i.stack.imgur.com/KdkBS.jpg)](https://i.stack.imgur.com/KdkBS.jpg)




**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](https://en.wikipedia.org/wiki/Multipole_expansion#Molecular_moments) 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](https://en.wikibooks.org/wiki/Molecular_Simulation/Quadrupole-Quadrupole_Interactions)


The problem is that most of these interactions die off very quickly. The quadrupole-quadrupole term is:[1](https://en.wikibooks.org/wiki/Molecular_Simulation/Quadrupole-Quadrupole_Interactions)


$$E(r)=\frac{-\Theta\_1\Theta\_2}{4\pi\epsilon\_0r^5}\times\Gamma(\theta\_1,\theta\_2,\phi)$$


So roughly $1/r^5$, compared to $1/r^3$ for dipole-dipole interactions, or $1/r^6$ for [dispersion forces](https://en.wikipedia.org/wiki/London_dispersion_force) 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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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](https://i.stack.imgur.com/5ASDO.jpg)](https://i.stack.imgur.com/5ASDO.jpg)


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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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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