Brook H. Rollins V, Jordan Davis, Evan Sova
8-16-12
Dr. Metzker
Chem 3010
Trends and Characteristics of the Periodic Table
The ionization energy of an atom is the energy required to remove electrons
to form cations. Which is in direct relation to the atomic radii and orbital
configuration. As the atomic radii decreases across a period, the ionization energy
increases. This is due to the atom’s smaller size, which pulls electrons closer to the
nucleus, making it harder to remove the electrons. But when electrons are added
(such as moving down a group or to an additional sub shell) the atomic radii
increases, making outer electrons easier to be removed. Exceptions are boron and
mercury, because removal of one of these outer electrons actually results in making
the atoms unstable and higher ionization energy. Beryllium and boron are some
exceptions to the general ionization trend. According to the trend boron should
technically have a higher ionization than that of beryllium. This case is caused by
beryllium having a full 2s subshell causing it to be stable. In boron the 2p electron is
easier to remove than the stable 2s in beryllium. Figure 1 shows the general trend
for ionization trend.
Figure 1-Periodic trend for Ionization Energy
Electron affinity is an amount of energy released by an atom when it
becomes ionized; it increases across a given period due to an increase in the valence
electrons. The trend of electron affinity is increasing from left to right across the
period and increasing from top to bottom down a group. Exceptions to these trends
generally are found in group 1A to group 2A and group 5A to group 6A. Both 1A and
5A have a higher electron affinity than their following group. This is caused by the
stability created by having subshells (s and p) that are half full. This causes the
energy of these groups to be higher than expected. The general trend is shown in
Figure 2.
Figure 2- Electron Affinity General Trend
Electronegativity is a measure of the ability of an atom to draw an electron to
itself. Electronegativity increases across a given period due to the increasing
effective nuclear charge. The effective nuclear charge is an attractive force between
the electrons in the electron cloud and its positively charged nucleus. While this
value increases, attracting electrons becomes more favorable for attachment of the
atom. The electronegativity trend also decreases down the periodic table due to
increasing atomic radii; therefore decreasing the effective nuclear charge and
making it less favorable for an atom to attract an electron. The exception to the
trend is Argon, because Argon is not believed to form bonds. The periodic trend is
shown in Figure 3.
Figure 3- Electronegativity Periodic Trend
The atomic radii is very easy to see and spot out on the periodic table, but
sometimes the empirical value is somewhat skewed. By definition of atomic radii, it
is the size of the isolated neutral atom and the length is measured in angstroms (Å).
The general trend for atomic radii found to be decreasing along the period and
increasing down each group. Exceptions to the trend are the noble gases. The noble
gases cannot be measured due to their extreme stability due to the octet being full.
Another exception is the transitional metals in groups 12-15. This is caused by the
electrons being added to the inner shell before the outer shell to increase shielding
of the outer shell and thus increasing stability. The general trend is show in Figure 4.
A graphic diagram is shown in Figure 5.
Figure 4- Periodic Trend for Atomic Radii
Figure 5- Graphic Trend for Atomic Radii
The melting point versus the atomic number trend depicts that with
increasing atomic number the melting points tend to increase. The reasoning for
why the melting point fluctuates by increasing and decreasing with the increasing
atomic numbers is due to the varying states of matter for each of the elements. For
example, the first two elements on the periodic table, hydrogen and helium are both
naturally occurring gases and the next few elements are metals, which significantly
increase in melting point values. The reason for why metals have higher melting
points are due to the high amount of energy needed to break bonds to change the
solid phase of a substance to a liquid phase. The stronger the bond between the
atoms of an element, the higher the energy required in breaking that particular
bond. Temperature in this case is the energy being used to break the bond to change
the solid to a liquid.
Figure 6- Periodic Trend for Melting Point