I submitted a question previously about why Nitrogen's atomic radius is lower
than both Oxygen's and Carbon's radius. I checked the archives and all I could
find out was that radii
tend to be larger as with the higher shells and from
right to left. What I don't understand is that from Boron to Neon this
pattern
holds good except at Nitrogen. Nitrogen's radius is smaller than both Carbon
and Oxygen. Why is Nitrogen an "oddball" compared to its neighbors on the
periodic table? Why doesn't its outer unpaired electron repell other electrons
thus being less stable and therefore having a higher radius?
First, check out the bottom of the answer I linked above. This gives a hint of
the explanation. Atomic radius tends to go with first ionization energy; the
higher the IE, the smaller the radius. This makes sense; IE is a measure of how
tightly electrons are held in the atom.
Now, how do we explain why nitrogen has a smaller radius/higher IE than either
oxgyen or carbon? The IE "hiccup" was explained, very briefly, at the
bottom of my
earlier answer. A more detailed explanation of what's going on follows
(with apologies to those who recognize that the atomic orbital model is only
approximately true for atoms other than hydrogen...)
- As electrons are added to atoms (new elements), protons are also added to
the nucleus, and each outer ("valence") electron feels a stronger
attraction from that nucleus because, on average, valence electrons are all the
same distance from the atom. Therefore, atomic radius decreases with
increasing atomic number.
- Second row elements, specifically carbon, nitrogen and oxygen, have four
subshells ("orbitals") in their valence shell: one s orbital
and three p orbitals. The s orbital is a bit lower in energy, and more
tightly held, than the p orbitals.
Hund's rule says that orbitals of equal energy are filled one electron at a
time. This may be explained by considering that two electrons in the same
orbital occupy the same specific region of space, and so repel each other more
than electrons in different orbitals. This is not a big energy difference, but
nevertheless we expect atomic radius to go up a little when this
happens.
- Let's consider what happens going from carbon, to nitrogen, to oxygen:
- In atomic carbon, there are four valence electrons: two in the lower-
energy s orbital, and one in each of two p orbitals.
- In atomic nitrogen, there are five valence electrons: two in the s orbital,
and one in each of the three p orbitals. However, since the p-electrons are
equally distant from the nucleus (they don't "shield" each other),
they each see a larger nuclear charge than in carbon, +7 rather than +6, and
the atomic radius drops.
- In atomic oxygen, there are six valence electrons: two in the s orbital,
one in each of three p orbitals, and one more which is jammed into one
of the p orbitals with another electron. The p electrons see a larger nuclear
charge (+8 rather than +7), and we would expect them to be more tightly held--
except that electron-electron repulsion is greater than in nitrogen,
and the net effect is that the atomic radius of oxygen is a bit more than that
of nitrogen, just as its first ionization energy is a bit less.
|