Atomic
Periodic Properties
The periodic variation in electron
configurations as one moves sequentially through
the Periodic Table from H to ever heavier elements
produces a periodic variation in a variety of
properties. We have
already seen the periodic
variation in atomic shape, and here we look at
three other properties. For more examples,
including these three, visit the
Dartmouth
College Periodic Puzzle in the ChemLab website.
The first such property we will consider is
atomic size as measured by the atomic
radius, a property discussed in the Zumdahl
text starting on page 561. (And remember: most
elements are found under ordinary conditions as
diatomic molecules—N2,
O2, etc.—polyatomic
molecules—P4, S8,
etc.—or solid metals. Here we are considering
only free atoms, often produced at high
temperatures only, and in the gas phase only.) The
figure below shows how atomic radius varies
qualitatively across the Periodic Table.
The general trend is easy to spot and to
understand: heavier atoms (those with large atomic
numbers) have more electrons, and increasing the
number of electrons means placing electrons into
orbitals with ever-increasing principal quantum
number—orbitals with ever-increasing size.
This clearly explains the trend in increasing size
as we go down any one column: H is
smaller than Li, etc., ending the first
column with the biggest atom for which we have
reliable data, Cs. These atoms' sizes are
governed by the size of their single highest-energy
electron's orbital, which is 1s for H
but 6s for Cs.
As we go across a row (look at Li
through Ne, for example), the size
decreases, but rather slowly: we are adding
electrons into orbitals with (for this row) the
same principal quantum number. These
orbitals are slowly shrinking in size because the
nuclear charge is increasing as we go across
a row, and increasing nuclear charge means
increasing the force attracting electrons to the
nucleus, making the orbitals contract.
Next, we look at two energetic properties of
atoms, the ionization energy (IE) and
the electron affinity (EA). The
ionization energy is the energy required to remove
the least tightly bound electron from an atom,
producing a positive ion and a free electron:
(When we start with a neutral atom, as we have
here, we say the energy change is the first
ionization energy. If we then remove another
electron, as in A+ —>
A2+ + e–, we call the energy need
to do so the second ionization energy, and
so forth, until we run out of electrons.) The
figure below shows the trend in first ionization
energies across the Periodic Table.
Compare this figure to the earlier figure of
Atomic Radii and note that in general, small
size means large ionization energy. This trend
is easy to understand: small atoms have few
electrons that are close to the nucleus. The closer
an electron is to a nucleus, the more energy it
takes to remove that electron from the atom. Now
look at the trends going across any one row of the
Periodic Table: ionization energies increase. For
example, the He ionization energy is greater than
the H ionization energy. Both have electrons in the
1s orbital, but the increased nuclear charge
of He pulls its two 1s electrons closer to
the nucleus, and thus more energy is needed to
remove either one of them than is needed for H with
only a single nuclear proton. Likewise, as we go
from Li to Ne, we are adding
electrons to the n = 2 orbitals (2s at
first, then 2p). These orbitals are pulled
closer to the nucleus as we go from Li to
Ne because the nuclear charge is increasing,
pulling the electrons closer. As we go down a
column (consider H through Cs or
He through Rn, for example), the
highest energy (outermost spatially) electron or
electrons will be found in orbitals of increasing
principal quantum number n: n = 1 for H and
He, 2 for Li and Ne, etc.,
through 6 for Cs and Rn. As the
principal quantum number of the electron we are
trying to remove increases, the energy needed to
remove it decreases even though the nuclear charge
is increasing as we go down a column. We explain
this through the concept of shielding: the
electrons in orbitals of smaller principal quantum
number shield (or screen) some of the nuclear
charge, causing our outer electron of interest to
think it is bound to an atom of roughly the same
nuclear charge no matter which row it is in.
Finally, we focus on the electron
affinity. This is the energy required to remove
the outermost electron from an atom that has one
extra electron stuck to it:
In other words, the electron affinity is the
ionization energy of the singly charged atomic
anion. The trends in electron affinities are
shown below.
Note first the elements that fall in the low end
of the range. For these elements, the electron
affinity is either zero or very close to zero,
which means that these elements do not form stable
anions. If you compare the locations of these
elements to the Periodic
Table of spherical elements that we already
discussed, you will see that those atoms that are
spherical because they have closed shells or
subshells also have zero or nearly zero electron
affinities. Likewise, those that are spherical
because they have half-filled shells or subshells
also have zero or very small electron affinities.
Why? Consider the rare gases, starting with
He. Adding an electron to He means
changing the electron configuration from
1s2 to
1s22s1, and the
screened He nuclear charge simply isn't
strong enough to bind that 2s electron and
make a stable He– anion. A similar
argument holds for all the other rare gases, and
for the alkaline earth elements, Be through
Ra, plus the elements at the end of the
transition metals, Zn, Cd, and
Hg, the extra electron would go into
a new subshell (as in Be
1s22s2 –>
Be–
1s22s22p1).
Again, the screened nuclear charge cannot bind such
an electron. For the half-filled elements such as
N
(1s22s22p3),
the three p electrons, you'll recall, are in
different p orbitals, one for each m
quantum number value. Adding a fourth to make
N– forces two electrons into the same
p orbital. Since electrons repel and since
the nuclear charge doesn't increase, that fourth
electron is not bound. N– does not
exist as a free atomic anion. (For elements below
N, you'll notice that the electron affinity
is not zero, but it is small. Adding that one extra
electron is not quite such a bad idea for these
elements because the electron finds itself in a
larger p orbital (3p for P,
4p for as, etc.), and these larger
orbitals allow the two electrons that are now
forced into the same orbital to avoid each other
better.)
But for the halogens, F through
At, the electron affinity is quite high.
These elements are just one electron away from a
closed shell configuration, and the extra electron
can be held by the attractive force of the atom's
nucleus. Note as well that Cu, Ag,
and Au, which are only one electron away
from a closed d subshell, also have high
electron affinities.
 9 Section Home
|
