Download Atomic Size (Atomic Radius)

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.