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Homework Help: About the dispersion force in polar molecules

  1. Jan 5, 2013 #1
    1. The problem statement, all variables and given/known data
    dispersion force exists in non-polar molecules due to instantaneous dipole.
    In polar molecule,the intermolecular force is the sum of dipole-dipole force and dispersion force.
    Polar molecules have permanent dipoles,this enables the oppositely charged end of molecules to attract each other.But when the effect of dipole-dipole force and dispersion force occurs simultaneously,I get confused.Why do dispersion forces exist in polar molecules?The positive end and the negative end of molecules are constantly attracted to each other.When molecules have permanent dipoles,why do we have to consider about the instantanous dipole?In my opinion,I think that dipole-dipole force is the "advanced version" of the dispersion force.The former exists in polar molecoles,while the latter only exists in non-polar

    Can anyone explain to me?I don't understand..
    Thx a lot!

    2. Relevant equations

    3. The attempt at a solution
  2. jcsd
  3. Jan 8, 2013 #2

    I too would love to hear an explanation for this.
  4. Jan 9, 2013 #3


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    Remember that the dispersion forces arise from electron-electron interactions. If two molecules have an electron cloud, you can be assured that the effect will be present. How much that effect affects the intermolecular picture overall varies by case but it is not correct to say that dispersion forces are not present in polar molecules.
  5. Jan 9, 2013 #4
    I think what he is asking is why would you account for both the dispersion forces AND dipole-dipole forces? Both situations explain different behaviors of electron densities. A molecule that only exhibits dispersion forces is very changeable and polarizable. However, a molecule exhibiting dipole-dipole forces has a clearly defined direction and magnitude. If you are defining the direction and magnitude, you have already explained the behavior of the electrons, and you wouldn't need to account for dispersion forces.

    But I think that this is incorrect reasoning. My guess would be that because all molecules have changing electron densities, they all exhibit random instantaneous dipoles. A dipole-dipole molecule may have generally defined and fairly permanent poles, but it is still changing and "wiggling." Take H2O for example. We know that oxygen is a fair amount more electronegative than hydrogen; individually, each of the H-O bonds' electrons are being shared unequally in favor of oxygen. Accounting for both bonds, we can predict an average pole. However, each individual bond between H and O should be accounted for. They can be said to point generally towards oxygen, but only on average. At any one point in time, we can not tell exactly where the pole is pointing, and so a decent amount of the time the polarization won't be perfectly towards oxygen. This is why we add them both; because both situations explain different behaviors of electron densities, they each need to be taken into consideration.

    In short, it doesn't matter that an individual molecule has a permanent dipole, because it's polarization can still be affected by neighboring molecules.
  6. Jan 10, 2013 #5
    Thx for your reply!Actually I have the same idea as you before I read your post :approve:,
    Anyway really thx :)
  7. Jan 10, 2013 #6


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    You always account for them. How much it matters varies by case but it is never insignificant. In the case of ammonia, which has a permanent dipole, London forces contribute more than half of all of the interaction forces. It isn't black or white. It's various shades of gray.
  8. Apr 3, 2014 #7
    I hate to revive a thread that is over a year old, but I feel that I must explain myself. Paragraph one was me explaining OP's original train of thought. Paragraph two was intended to walk OP through my reasoning steps as to why that [sentence you bolded] did not make sense.
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