My teacher told me that Neon has a larger atomic radius than Fluorine.I am of the understanding that it is merely a consequence of the way we define the atomic radius and that we use Van der walls radius for neon which is larger that the covalent radius of fluorine.

He also said that the radius of a fluorine 1- anion is more than a neon atom by using the fact that they are iso-electronic species and that since neon has a higher nuclear charge, it will pull on the 10 electrons more than fluorine.

But what definition of atomic radius am i using for neon now? Aren't we always supposed to use the Van der walls definition for noble gases? In case we do use the Van der walls radius of the neon atom and the ionic radius of the fluorine anion, how am i supposed to compare these two and come to the conclusion that the fluorine anion is larger?


1 Answer 1


Atomic radii cannot really be uniquely nor accurately defined. However, there is generally a very big difference between covalent, van der Waals, and ionic radii.

So, by most any definition of the respective terms the covalent radius of fluorine will be smaller than the van der Waals radius of neon. Similarly, the ionic radius of fluoride ion will be larger than the van der Waals radius of neon. The reason for the latter is the higher nuclear charge on neon.

However, most sources list a slightly smaller van der Waals radius for fluorine compared to neon, but a slightly larger covalent radius for fluorine compared neon. So when comparing radii you have to carefully specify what kind of radii you are talking about.

  • $\begingroup$ So we can generalize it to say that the atomic radius of fluoride is greater than neon (as long as we aren't comparing van der Waals radius.) ? $\endgroup$
    – SMcCK
    Jun 2, 2015 at 10:08

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