Chemistry
Periodic Trends
E. Spinelli
There are four principle Periodic Trends:
I.
II.
III
IV
Particle Size (Atomic and Ionic Radii)
Ionization Energy
Electron Affinity
Electronegativity
I. Particle Sizes are generally defined as the measurement of the distance from the nucleus to the outermost electrons.
Since the outermost portion of an atom or ion consists of an electron cloud these sizes are considered close
approximations. Particle size is thought of as the volume that contains 90 percent of the electron density of the atom.
The sizes of atoms of the elements depend on the state in which we may find them. They can generally be found in one
of three possible states: 1) a free uncombined state as monatomic atoms, 2) ionicly or covalently chemically bonded to
another atom, or 3) uncombined or bonded but are under the influence of an attraction from another atom or molecule
(van der Waals radii).

Atomic Radii are determined by taking half of the measured distance between adjacent nuclei of two identical
metal atoms.

Ionic Radii are determined by taking half of the measured distance between adjacent nuclei of two identical ions
that are found within an ionicly bonded compound.

Covalent Radii are determined by taking half of the measured distance between adjacent nuclei of two identical
atoms that are found within a covalently bonded molecule.

van der Waals Radii are determined by taking half of the measured distance of closest approach between
adjacent nuclei of two identical atoms that are not bonded together.
Fundamental forces that determine the size of particles:

Electrostatic attraction of the positively charged nucleus for the negatively charge electrons.

Repulsive forces between the electrons within the orbitals of the electron clouds.
Atomic Radius Trend for the Representative Elements:
Atomic radius increases from top to bottom within a group.
Explanation:
1. Outermost electrons (valence electrons) are farther from the nucleus due to a higher principal quantum number
and therefore the electron’s attraction to the nucleus (effective nuclear charge) is not as strong resulting in a
larger atom.
2. Lower elements (higher atomic number) in a group have more energy levels than those above. Energy levels
between the nucleus and the valence electrons reduce the electrostatic attraction between the protons and the
electrons (Shielding Effect). The lower attraction causes the electrons to not to be drawn in as closely and
therefore, results in a larger atom.
Atomic radius decreases from left to right across a period.
Explanation:
As atomic number increase from left to right across a period the nuclear charge increases while the additional
electrons are completing orbitals with the same energy level and therefore, are not any farther from the nucleus.
This lack of the electron’s ability to produce any shielding effect causes a net increase in effective nuclear
charge and result in a decrease in atomic radius.
Ionic Radius Trends for the Representative Elements
Cations are smaller than their respective neutral atoms.
Explanation:
Fewer electrons than protons cause an increase in effective nuclear charge and consequently a decrease in
repulsion between the electrons resulting in a smaller ion.
Anions are larger than their respective neutral atoms.
Explanation:
Greater electrons than protons cause a decrease in effective nuclear charge and consequently an increase in
repulsion between the electrons resulting in a larger ion.
Cations and anions lower within a group are larger than those above.
Explanation:
Cations and anions that are lower have more energy levels and consequently have a greater shielding effect
which causes a lower repulsion between electrons and therefore are larger ions than those above them.
Cations that have a greater positive charge are smaller than those with a lower positive charge.
Explanation:
The greater positive charge (fewer electrons) represents a greater effective nuclear charge and therefore a
smaller ion.
Anions that have a greater negative charge are larger than those with a lower negative charge.
Explanation:
The greater negative charge (more electrons) represents a lower effective nuclear charge and therefore a
larger ion.
II. Ionization Energy is the minimum energy required to remove an electron from a gaseous atom in its ground state.
Generally, the greater the electrostatic attraction of the nucleus for the electrons the greater the ionization energy.
Fundamental forces that determine the ionization energy of an atom.

Particle size; the smaller the particle the greater the attraction for the electrons.

Shielding Effect; the fewer the energy levels between the nucleus and the valence electrons the greater the
attraction for the electrons.

Effective Nuclear Charge; the fewer the number of electrons than protons the greater the attraction for the
electrons.

Orbital Configuration; electron configurations that are completely filled or to a lesser degree, half-filled,
represent more stable arrangements that have lower potential energy and therefore require greater ionization
energy.
Atoms of elements higher within a group require greater ionization energy than those that are lower.
Explanation:
Elements higher in a group are smaller in radius and have fewer energy levels between the nucleus and valence
electrons (lower shielding effect) resulting in a greater attraction for the electrons and consequently require
greater ionization energy.
Atoms of elements farther to right within a period require greater ionization energy than those to the left.
Explanation:
Except for a few notable irregularities* (explained later), ionization energy increases across a period with an
increase in atomic number. The increase is due to the increase in effective nuclear charge. While the number of
electrons is also increasing, these electrons are added to the same energy level that is not any farther from the
nucleus (as in the trend for decrease in atomic radius, explained above). Therefore the increased nuclear charge
results in a stronger attraction thus requiring greater ionization energy to remove an electron.
Noble gases have the highest ionization energy of their respective periods.
Explanation:
Noble gases have completely filled valance shells and therefore have the most stable ground state electron
configurations. These stable configurations represent especially low potential kinetic energy thus requiring
high ionization energy to remove an electron.
*Irregularities to the general trend in left to right increase in ionization energy across a period.
Explanation:
There are two notable exceptions to the general trend explained above. The first occurs between group II and
group III of the representative elements. Contrary to the general trend, group III elements actually have a lower
first ionization energy than group II elements. This is due to group III having a single p orbital electron that is
shielded not only by its inner electrons but also by its filled valence s orbital. This shielding effect of the s orbital
electrons results in a lower ionization energy. A second irregularity to the general trend occurs between group V
and group VI of the representative elements. Again contrary to the general trend, group VI actually has a lower
first ionization energy than group V elements. The three p electrons in group V elements are distributed among
separate orbitals. However, the fourth p electron of group VI elements is paired with one of the other three p
electrons in a single orbital. This pairing of electrons in a single orbital causes an increased electrostatic
repulsion between the electrons and results in a decrease in ionization energy.
Subsequent ionizations always require greater ionization energy.
Explanation:
The loss of an electron from an atom results in an increase in effective nuclear charge and a decrease in
electrostatic repulsion among the remaining electrons. The more electrons removed the higher the effective
nuclear charge, the lower the repulsion, and consequently the higher the ionization energy required. Additional
electrons removed beyond the valence shell require especially large quantities of ionization energy due not only
to the increased nuclear charge and decreased repulsive forces but also due to the effect of the stability of the
configurations of the inner core electrons that have particularly low potential kinetic energies. In fact, an atom’s
electrons can be predicted from an analysis of its successive ionizations as illustrated in the figure below.
III. Electron Affinity is the change in energy that occurs when an atom in the gaseous state acquires an electron to form
an anion. Elements that achieve more stable configurations will have lower potential kinetic energies and consequently
emit the difference in energy as a negative value typical of exothermic reactions. However, the energy released is often
thought of in positive terms therefore the sign of the energy value is changed. Consequently, those elements that have a
tendency to acquire an electron will be reported as having relatively large positive values while those elements that result
with higher potential kinetic energies after having an electron incorporated into their configurations will be reported will low
positive, near zero, or even negative values for electron affinity.
Atoms higher within a group have higher electron affinities than those below.
Explanation:
Similar to the trend for ionization energy, higher electron affinities are caused by the elements having smaller radii
and fewer energy levels between the nucleus and valence electrons (lower shielding effect) resulting in a greater
attraction for their existing electrons and also any additional electron.
Atoms of elements farther to right within a period have greater electron affinity than those to the left.
Explanation:
Again similar to ionization energy and with the same few notable irregularities* (explained later), electron affinity
increases across a period with an increase in atomic number. The increase is due to the increase in effective
nuclear charge. While the number of electrons is also increasing, these electrons are added to the same energy
level that is not any farther from the nucleus (as in the trend for decrease in atomic radius, explained above).
Therefore the increased nuclear charge results in a stronger attraction thus producing greater electron affinity
when an electron is acquired.
Noble gases have especially low electron affinities.
Explanation:
Noble gases have completely filled valance shells and therefore have the most stable ground state electron
configurations. These stable configurations represent especially low potential kinetic energy thus have very
little tendency to acquire an additional electron.
*Irregularities to the general trend in left to right increase in ionization energy across a period.
Explanation:
As with ionization energy, there are two notable exceptions to the general trend explained above. The first occurs
between group II and group III of the representative elements. Contrary to the general trend, group II elements
actually have a lower electron affinity than group II elements. This is due to group III having a single p orbital
electron that is shielded not only by its inner electrons but also by its filled valence s orbital. This shielding effect
of the s orbital electrons results in a lower electron affinity. A second irregularity to the general trend occurs
between group V and group VI of the representative elements. Again contrary to the general trend, group VI
actually has a lower electron affinity than group V elements. The three p electrons in group V elements are
distributed among separate orbitals. However, the fourth p electron of group VI elements is paired with one of the
other three p electrons in a single orbital. This pairing of electrons in a single orbital causes an increased
electrostatic repulsion between the electrons and results in a decrease in electron affinity.
Note: Although an oxygen atom has a positive electron affinity, indicating that the acquired electron produces a more
stable configuration and has a lower potential kinetic energy, the electron affinity of the O1- ion has a large negative value,
thus indicating it is less stable and has a higher potential kinetic energy despite the O2- ion is isoelectronic with the noble
gas neon. This is true because in the gas phase the O2- ion’s increase in electrons produces an electron repulsion that
outweighs the increase in stability of its configuration. However, in the solid state the O2- ion’s configuration and
consequent lowering of its potential kinetic energy is achieved due to electrostatic attractions of surrounding cations.
IV Electronegativity is the relative tendency of an atom to attract toward itself the electrons in a chemical bond.
Electronegativity can not be measured directly. It is a relative concept whose value is determined compared to other
elements. It is particularly helpful in distinguishing bond character. If the electronegativity difference between two
elements is equal to or greater than 1.67 then the bond is considered to be principally ionic. Electronegativity differences
greater than zero but less than 1.67, are considered to be principally polar covalent. Add finally, electronegativity
differences close to zero are considered to be principally nonpolar covalent.
Electronegativity follows a pattern similar to that of ionization energy and electron affinity, and for similar reasons. The
exception for the noble gases is explained by the stability of the elements’ electron configuration and consequent low
potential kinetic energies. Krypton and Xenon however, are large enough atoms with significant shielding effect to off set
the repulsive forces of electrons observed in smaller atoms. The value for Radon has not been determined primarily due
to its radioactivity.