Why NCl3 is dipole-dipole in intermolecular force?

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Discussion Overview

The discussion centers on the intermolecular forces present in nitrogen trichloride (NCl3), specifically questioning why it exhibits dipole-dipole interactions rather than London dispersion forces. Participants explore the concepts of electronegativity, bond polarity, and the nature of molecular dipoles.

Discussion Character

  • Debate/contested
  • Conceptual clarification
  • Technical explanation

Main Points Raised

  • One participant questions why NCl3 is classified as dipole-dipole given that nitrogen and chlorine have the same rounded electronegativity value of 3.0, suggesting it should exhibit London dispersion forces instead.
  • Another participant argues that while the electronegativities are similar, the small difference is significant enough to result in dipole-dipole interactions, as it outweighs the weaker van-der-Waals forces.
  • A participant references a source indicating that polar bonds occur with an electronegativity difference between 0.4 and 1.7, noting that the difference of 0.2 in NCl3 is insufficient for a polar bond.
  • Some participants emphasize that there is no clear boundary between polar and nonpolar bonds, suggesting that the degree of polarity can vary and is not strictly defined.
  • One participant asserts that NCl3 should have a permanent dipole moment due to the asymmetrical charge distribution caused by the lone pair on nitrogen, supported by experimental measurements of its dipole moment.
  • Another participant suggests that while every bond between different types of atoms is somewhat polar, determining whether a bond is polar can be subjective and may rely on arbitrary thresholds.

Areas of Agreement / Disagreement

Participants express differing views on the nature of bond polarity and the classification of NCl3's intermolecular forces. There is no consensus on the criteria for defining polar bonds or the implications of electronegativity differences.

Contextual Notes

Participants note the lack of a definitive threshold for bond polarity and the subjective nature of determining whether a bond is polar, highlighting the complexity of the topic.

terryds
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3.0 is a rounded value, they are not exactly the same. The small difference is still more important than the weaker van-der-Waals forces.
This post suggests a notable difference between the electronegativity.
 
mfb said:
3.0 is a rounded value, they are not exactly the same. The small difference is still more important than the weaker van-der-Waals forces.
This post suggests a notable difference between the electronegativity.
But, according to https://en.wikipedia.org/wiki/Chemical_polarity ,
  • Polar bonds occur when the difference in electronegativity between the two atoms is between 0.4 and 1.7
In that post you suggested, the electronegativity difference is just 0.2 , still not enough to perform a polar bond..
 
There is no sharp line dividing polar and nonpolar. Some bonds are more polar than others.
In the same way, there is no sharp line dividing polar and ionic bonds.
 
mfb said:
There is no sharp line dividing polar and nonpolar. Some bonds are more polar than others.
In the same way, there is no sharp line dividing polar and ionic bonds.

If there is no sharp line dividing polar and non-polar, how to determine if such bond is polar?
By memorizing only??
 
I don't see no reason why NCl3, even with perfectly covalent bonds, should not have a permanent dipole moment (which we know from experimental measurements to be 0.6 D). It is a question of the asymmetry of the electronic charge distribution around the nuclei. The lone pair on N is especially important in this respect.
 
Last edited:
terryds said:
If there is no sharp line dividing polar and non-polar, how to determine if such bond is polar?
By memorizing only??
Every bond between atoms of different types is a bit polar. You can set an arbitrary threshold for saying "this is sufficient to call it polar", which can be convenient sometimes, but you do not have to.
 

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