Periodic Trends
In the SparkNote on the Periodic table we discussed a number of simple periodic trends. In this section we
will discuss a number of more complex trends, the understanding of which relies on knowledge of atomic
structure.
Before getting into these trends, we should engage a quick review and establish some terminology. As seen
in the previous section on the octet rule, atoms tend to lose or gain electrons in order to attain a full valence
shell and the stability a full valence shell imparts. Because electrons are negatively charged, an atom
becomes positively or negatively charged as it loses or gains an electron, respectively. Any atom or group
of atoms with a net charge (whether positive or negative) is called an ion. A positively charged ion is a
cation while a negatively charged ion is an anion.
Now we are ready to discuss the periodic trends of atomic size, ionization energy, electron affinity, and
electronnegativity.
Atomic Size (Atomic Radius)
The atomic size of an atom, also called the atomic radius, refers to the distance between an atom's nucleus
and its valence electrons. Remember, the closer an electron is to the nucleus, the lower its energy and the
more tightly it is held.
Moving Across a Period
Moving from left to right across a period, the atomic radius decreases. The nucleus of the atom gains
protons moving from left to right, increasing the positive charge of the nucleus and increasing the attractive
force of the nucleus upon the electrons. True, electrons are also added as the elements move from left to
right across a period, but these electrons reside in the same energy shell and do not offer increased
shielding.
Moving Down a Group
The atomic radius increases moving down a group. Once again protons are added moving down a group,
but so are new energy shells of electrons. The new energy shells provide shielding, allowing the valence
electrons to experience only a minimal amount of the protons' positive charge.
Cations and Anions
Cations and anions do not actually represent a periodic trend in terms of atomic radius, but they do affect
atomic radius, and so we will discuss them here.
A cation is positively charged, meaning that it is an atom that has lost an electron or electrons. The positive
charge of the nucleus is thus distributed over a smaller number of electrons and electron-electron repulsion
is decreased, meaning that the electrons are held more tightly and the atomic radius is smaller than in the
normal neutral atom. Anions, conversely, are negatively charged ions: atoms that have gained electrons. In
anions, electron-electron repulsion increases and the positive charge of the nucleus is distributed over a
large number of electrons. Anions have a greater atomic radius than the neutral atom from which they
derive.
Ionization Energy and Electron Affinity
The process of gaining or losing an electron requires energy. There are two common ways to measure this
energy change: ionization energy and electron affinity.
Ionization Energy
The ionization energy is the energy it takes to fully remove an electron from the atom. When several
electrons are removed from an atom, the energy that it takes to remove the first electron is called the first
ionization energy, the energy it takes to remove the second electron is the second ionization energy, and so
on. In general, the second ionization energy is greater than first ionization energy. This is because the first
electron removed feels the effect of shielding by the second electron and is therefore less strongly attracted
to the nucleus. If a particular ionization energy follows a previous electron loss that emptied a subshell, the
next ionization energy will take a rather large leap, rather than follow its normal gently increasing trend.
This fact helps to show that just as electrons are more stable when they have a full valence shell, they are
also relatively more stable when they at least have a full subshell.
Ionization Energy Across a Period
Ionization energy predictably increases moving across the periodic table from left to right. Just as we
described in the case of atomic size, moving from left to right, the number of protons increases. The
electrons also increase in number, but without adding new shells or shielding. From left to right, the
electrons therefore become more tightly held meaning it takes more energy to pry them loose. This fact
gives a physical basis to the octet rule, which states that elements with few valence electrons (those on the
left of the periodic table) readily give those electrons up in order to attain a full octet within their inner
shells, while those with many valence electrons tend to gain electrons. The electrons on the left tend to lose
electrons since their ionization energy is so low (it takes such little energy to remove an electron) while
those on the right tend to gain electrons since their nucleus has a powerful positive force and their
ionization energy is high. Note that ionization energy does show a sensitivity to the filling of subshells; in
moving from group 12 to group 13 for example, after the d shell has been filled, ionization energy actually
drops. In general, though, the trend is of increasing ionziation energy from left to right.
Ionization Energy Down a Group
Ionization energy decreases moving down a group for the same reason atomic size increases: electrons add
new shells creating extra shielding that supersedes the addition of protons. The atomic radius increases, as
does the energy of the valence electrons. This means it takes less energy to remove an electron, which is
what ionization energy measures.
Electron Affinity
An atom's electron affinity is the energy change in an atom when that atom gains an electron. The sign of
the electron affinity can be confusing. When an atom gains an electron and becomes more stable, its
potential energy decreases: upon gaining an electron the atom gives off energy and the electron affinity is
negative. When an atom becomes less stable upon gaining an electron, its potential energy increases, which
implies that the atom gains energy as it acquires the electron. In such a case, the atom's electron affinity is
positive. An atom with a negative electron affinity is far more likely to gain electrons.
Electron Affinities Across a Period
Electron affinities becoming increasingly negative from left to right. Just as in ionization energy, this trend
conforms to and helps explain the octet rule. The octet rule states that atoms with close to full valence
shells will tend to gain electrons. Such atoms are located on the right of the periodic table and have very
negative electron affinities, meaning they give off a great deal of energy upon gaining an electron and
become more stable. Be careful, though: the nobel gases, located in the extreme right hand column of the
periodic table do not conform to this trend. Noble gases have full valence shells, are very stable, and do not
want to add more electrons: noble gas electron affinities are positive. Similarly, atoms with full subshells
also have more positive electron affinities (are less attractive of electrons) than the elements around them.
Electron Affinities Down a Group
Electron affinities change little moving down a group, though they do generally become slightly more
positive (less attractive toward electrons). The biggest exception to this rule are the third period elements,
which often have more negative electron affinities than the corresponding elements in the second period.
For this reason, Chlorine, Cl, (group VIIa and period 3) has the most negative electron affinity.
Electronegativity
Electronegativity refers to the ability of an atom to attract the electrons of another atom to it when those
two atoms are associated through a bond. Electronegativity is based on an atom's ionization energy and
electron affinity. For that reason, electronegativity follows similar trends as its two constituent measures.
Electronegativity generally increases moving across a period and decreases moving down a group. Flourine
(F), in group VIIa and period 2, is the most powerfully electronegative of the elements. Electronegativity
plays a very large role in the processes of Chemical Bonding.