Kelly Taylor
Belinda Schaafsma
Ben Ottoson
Ben Gibson
Periodic Trends Report
Electronegativity
Electronegativity is the ability that an atom in a molecule has to attract shared electrons
to itself. This trend is connected to atomic number and distance that valence electrons are from
the charged nucleus which can be seen below in Figure 1. Electronegativity, in general, from left
to right across the period table increases because the electrons are strongly attracted to the
nucleus, it results in a stronger nuclear charger pulling on the shared electrons. Electronegativity
tends to decrease going down a group due to the increase in distance from the valence electrons
to the charged nucleus (i.e. a decreasing Zeff). Exceptions include transition metals as well as
noble gases (see figure 1). These exceptions will be discussed in more detail along with the other
exceptions for electron affinity, atomic radii, and ionization energy.
Figure 1: Pauling Electronegativities
Electron Affinity
Electron affinity is the change in energy of an atom when an electron is added. It can be
affected by nuclear charge, distance, and the shielding effect. In general, from left to right across
the periodic table, the electron affinity increases. From top to bottom, the electron affinity
decreases. This trend is related to the atomic radii which can be seen below in Figure 3. The
atomic radii decrease from left to right across the table making it easier to grab an extra atom
(i.e. higher electron affinity). Atomic radii increase from top to bottom lessening the effective
nuclear charge (Zeff) making it harder for the atom to grab an electron (i.e. lower electron
affinity).
Figure 2: Electron Affinities
Atomic Radii
Atomic radii is defined by calculating half of the distance between the nuclei in a
molecule containing identical atoms. The scale for this method is the most common method used
and is measured in Angstroms. There are two general trends involving atomic radii. The radii
decrease as we move across each period from left to right and increases as we move down each
group of the periodic table. Effective nuclear charge increases as we move left to right, which is
a result of poor shielding. The lack of shielding pulls the valence electrons closer to the nucleus
and thus, the size of the atom decreases. Also, as we move down each group, the increase in
orbital shells increases the size of the atom. Exceptions to these general trends include the noble
gases, which cannot be measured because the octet is completely stable and unable to form stable
bonds, as well as the transition metals from group 12-15. The exceptions for atomic radii trends
can be better seen in Figure 3 below and can be further explained later in this report.
Figure 3: Atomic Radii
Ionization Energy
Ionization energy is the amount of energy required to remove the highest energy electron
from an atom. One can see from Figure 4 below, the ionization energy increases from left to
right across the periodic table due to the decreasing atomic radii of the atoms. The smaller the
atom, the closer the electron is to the nuclei; this allows for an increasing Zeff, making it harder to
remove the electron. The ionization energy also decreases as you move from top to bottom on
the periodic table. This is a direct result of an increasing number in electron shells. This
increase in shells increases the shielding effect and decreases the Zeff.
Figure 4: Ionization Energy vs. Atomic Number
Exceptions to Periodic Trends
The d-block contraction gives evidence to the exception of transition metals for the
atomic radii trend, as seen in Figure 3. Transition metals have atomic radii that stay the same
which does not follow the trend. This is due to the fact of electrons being added to the inner dshell. The d electrons are poor at shielding when compared to s and p electrons. Other exceptions
to the electronegativity trend are noble gases because they have a full outer shell which does not
allow for space in order to attract shared electrons. An exception to ionization energy and
electron affinity trends is the decrease in electronegativity from Beryllium to Boron. Beryllium is
more stable than boron because it has filled subshells 1S and 2S. Boron is more likely to give up
its electron because it has one more subshell than Beryllium with only one electron. Another
exception to these two trends is the decrease in electronegativity from Nitrogen to Oxygen.
Nitrogen is more stable than Oxygen because, similar to Be and B, it has a half filled subshell
2P. Therefore, Oxygen has one more electron than a half filled shell. The repulsion between the
paired electrons makes it easier to remove one.
Advanced Data of Periodic Trends
Melting point is when a compound changes from a solid to a liquid. As the atomic
number increases across the periodic table from left to right the melting point increases and when
the atomic number increases down the periodic table the melting point decreases. This can be
related to the change in atomic radii. When the atomic radii increases, it is more reactive which
makes bonds easier to break. This can be due to the shielding effect of inner electrons in a large
multi-electron atom canceling out part of the nuclear charge. This can be shown through the
comparison of lithium and francium. Lithium has a melting point around 450 K while francium`s
melting point is around 300 K (see Figure 5)
There are exceptions to the general trends due to certain characteristics of varying atoms.
For example, carbon has the higher melting point than nitrogen or oxygen (Figure 5). This is due
to carbon being solid under natural conditions and being able to form stable bonds while nitrogen
and oxygen are naturally occurring gases.
Figure 5: Periodic Trend of Melting Point vs. Atomic Number
Lanthanide Contraction
The lanthanide contraction accounts for a greater decrease in atomic radii than expected
for elements in the lanthanide series. As shown in Figure 6, the atomic radii of these elements
decreases as the atomic number increases due to poor shielding affects of the 4f electrons, which
are weaker than the 3d electrons pertaining to shielding. This poor shielding can be related to the
5s and 5p orbitals penetrating the 4f sub-shell (see figure 7). This allows the increasing nuclear
charge to be more effective on the 6s electrons leading to the smaller than expected atomic radii.
Figure 6: Atomic Radius vs. Atomic Number
Figure 7- F orbital