Anion:
An Anion is an atom or group of atoms that contains more electrons than protons.
Because of the imbalance of charge the system will have a net charge that is
negative. Consequently, any negatively charged system is referred to as an anion.
Atomic Radius:
The Atomic Radius is a very common method of measuring the relative sizes of
atoms. It is determined as half the distance between two identical atoms when
bonded together. For instance, half the distance between to atoms of Oxygen in O2
will be the atomic radius of an Oxygen atom.
Cation:
A Cation is an atom or group of atoms that contains more protons than electrons.
Because of the imbalance of charge the system will have a net charge that is positive.
Consequently, any positively charged system is referred to as a cation.
Electron Affinity:
Electron Affinity is a measure of the desire or ability of an atom to gain electrons. It
is an energy concept. The formal definition states that Electron Affinity is the
amount of energy released when an electron as added to an atom. Most atoms tend
to lose energy when they gain electrons. Some atoms do not. Those that do not tend
to appear in the lower left corner of the Periodic Chart. The concept of Electron
Affinity tends to be viewed as an exothermic process. The elements located in the
upper right corner of the Periodic Chart have the high E.A. values while those in the
lower left corner have the low E.A. values. This is interpreted as meaning that
elements in the upper right corner are usually found an anions while those in the
lower left corner are usually not found as anions. A generic equation of the E.A.
process would be as follows.
Electronegativity:
The concept of Electronegativity refers to the ability of a bonded atom to draw
electrons towards itself. It is defined as the relative ability of an atom in a molecule
to attract electrons towards itself. As atoms bond, electrons are shared or
transferred. The decision about sharing or transferring electrons is made based
upon the electronegativities of the two bonded atoms. The atom with the higher
electronegativity will dominate the electrons. The greater the difference between the
electronegativity values of the two bonded atoms, the more the electrons will be
transferred and the less they will be shared. In order to be able to determine
electronegativity values it is important to observe the behavior of atoms in a bonded
situation. Consequently, the Noble Gases do not usually appear with listed
electronegativity values.
There are many different methods available for assigning electronegativity values.
Of all the options available, the most widely used set of values was determined by
Pauling and is called the Pauling Scale. It assigns the largest electronegativity value
to Fluorine and lowest to Francium. Fluorine has an electronegativity of 4.0, and
Francium has an electronegativity of 0.7. The numbers are unitless because they are
ratios.
Ionization Potential:
Ionization Potential is an energy term. It refers to the process of cation formation.
The definition states that Ionization Potential is the amount of energy required to
remove the highest energy electron from the valence level of an atom. Because
energy must be put into the system, the process in endothermic. Consequently, from
an energy point of view, this is an unfavorable process. Atoms in the upper right
hand corner of the Periodic Chart have the largest I.P. values and the elements in
the lower left corner of the Periodic Chart have the lowest I.P. values. Because the
elements in the lower left corner need the smaller amounts of energy to form
cations, these elements are usually found as cations in chemical process. The
elements in the upper right corner, on the other hand, need larger amounts of
energy to form cations. Therefore, they are not usually going to appear as cations.
The generic equation of the Ionization Potential process is as follows.
There are additional Ionization Potential terms, such as 2nd I.P., 3rd I.P., and so on.
This refers to the sequential removal of electrons from an atom. There will be one
I.P. value for each electron that at system has.
Noble Gas Configuration:
The Noble Gas Configuration is the preferred configuration for the arrangement of
electrons on all atoms. This configuration imparts a desirable level of stability into a
system. Consequently, all atoms will strive to gain or lose electrons in order to create
an electronic configuration that is isoelectronic with that of a Noble Gas. Usually,
the smaller atomic number elements are able to do this quite easily. The larger
atomic number elements are more restricted in their abilities to achieve this
condition. Therefore, many of the larger elements must resort to alternate methods
of achieving stability.
Oxidation Number:
The Oxidation Number of an atom or group of atoms is the charge that it carries.
The term is replacing the older terms of "charge" and "valence."
Pauling Scale:
The Pauling Scale is the most commonly used scale of electronegativity values. The
calculations used to arrive at the numbers in the scale are complex. It is most
common to simply know the results of those calculations. One of the major benefits
of the scale is the simplicity of the numbers. The scale is based on Fluorine having
the largest electronegativity with a value of 4.0. The Francium atom is assigned the
lowest electronegativity value at 0.7. All other values are located between these
extremes. As anchor points, the elements in the second period on the Periodic Chart
have values usually rounded off as follows:
Li--1.0 Be--1.5 B--2.0 C--2.5 N--3.0 O--3.5 F--4.0.
Periodic Chart Trends:
The Periodic Chart Trends refer to overall pattern of changes in Properties of
Atoms that take place within Groups and Periods on the Periodic Chart.
Frequently, it is more useful to know how different elements compare than it is to
know specific numerical information. Therefore, the Trends have been developed to
aid in doing comparative work. While there will be exceptions, or deviations, in the
Trends, the patterns are generally reliable.
The Trends are:

Within a Period from left-to-right:
 I.P. increases
 E.A. increases
 Electronegativity increases
 Size (radius) decreases

Within a Group from top-to-bottom:
 I.P. decreases
 E.A. decrease
 Electronegativity decreases
 Size (radius) increases
Radius:
The Radius of an atom is a difficult thing to measure. Because of the variable nature
of the electrons, the outside edge of an atom is poorly defined. Therefore, the outside
edge is not a reliable location for measuring to. As a result, the sizes of atoms are
measured in a variety of ways. Among them are terms such as the Atomic Radius
and the Covalent Radius. Each will have its own specific definition. Methods such as
this will provide relative sizes of atoms that are useful for comparison purposes.
Another option is to calculate the Radius of an atom, based on some arbitrary
standard. This will not provide an exact size for the atom, but will, again, provide
information that is useful for comparisons. This approach calculates the distance
between the center of the nucleus and the most probable position of the valence
level. The equation used for this purpose is
In this equation, n is the value of the valence level. Zeff is the effective nuclear charge
on a valence level electron. The term ao is a constant that represents the selection of
units for the radius. For instance, if ao is set at 0.54, then the resulting radius is
measured in Angstroms.
Stability Factors:
The Stability Factors referred to here are used to determine the numbers of
electrons that are gained or lost in chemical bonding processes. As atoms struggle to
improve stability by bonding, there are certain methods available to them.
Generally, there are four factors that appear most frequently. Listed here, in order
of decreasing importance:




Desire to achieve a configuration that is isoelectronic
with a Noble Gas.
Desire to lose all valence level electrons.
Desire to lose the valence level p orbital electrons.
Desire to maintain filled, empty, and half-filled sets of
orbitals.
The desires to lose valence level electrons or valence level p orbital electrons are
very similar in terms of preference. In some atoms one of the factors is more
important, and for other elements the other factor is more important. Regardless,
the idea of half-filling a set of orbitals will provide some stability, but it is definitely
a very low priority.