Section 3

Electron Configuration and Periodic Properties

Atomic Radii

Atomic Radius- Defined as one-half the distance between the nuclei of identical atoms that are bonded together.

Period Trends- The trend to smaller atoms across a period is caused by the increasing positive charge of the nucleus.

Group Trends- The atomic radii of the main-group elements increase down a group.

Ionization Energy

A + energy ---> A+ + e-

The A+ represents an ion of element A with a single positive charge.

Ion- is an atom or group of bonded atoms that has a positive or negative charge.

Ionization- The energy required to remove one electron from a neutral atom of an element.

Period Trends- Ionization energies of the main-group elements increase across each period.

Group Trends- Among main-group elements, ionization energies generally decrease down the groups. Electrons removed from atoms of each succeeding element in a group are in higher energy levels, farther from the nucleus. They are therefore removed more easily. As atomic number increases down a group, more electrons lie between the nucleus and the electrons in the highest occupied energy levels.

Removing Electrons from Positive Ions

If you have enough energy, electrons can be taken away from positive ions as well as the neutral ion. Each successive electron that is removed from an ion, feels an increasingly stronger effective nuclear charge.

Electron Affinity

The energy change that occurs when an electron is acquired by a neutral atom is called the atom’s electron affinity. Most atoms release energy when acquiring an electron.

A + e- ----> A- + energy

Some atoms are forced to gain an electron.

A + e- + energy -----> A-


Halogens have large negative electron affinities because they gain electrons most readily

Group Trends

Trends for electron affinities within groups are not as regular as the trends in ionization energies. In general, electrons add with greater difficulty down the group. This is because of a slight increase in effective nuclear charge down a group, which increases electron affinity. Also, an increase in atomic radius size down a group, which decreases electron affinity. The size effect is more prominent/over powering.

Adding Electrons to Negative Ions

It is more difficult to add an electron to an already negatively charged ion. Second electron affinities are always positive.

Ionic Radii

Cation- positive ion

Anion- negative ion

Period Trends-

The metals to the left tend to form cations and the nonmetals to the upper right tend to form anions

Cationic radii decrease across a period

Group Trends-

Gradual increase of atomic radii and ionic radii down a group


Valence electrons are the electrons available to be lost, gained, or shared in the formation of chemical compounds


Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound

Period Trends-

Electronegativities tend to increase across each period, except for the fact that some Noble gases do not form compounds, so their electronegativity can’t be estimated.

Group Trends-

Electronegativities either decrease down a group or remain about the same


Atomic Radii


Atomic Radii decrease across the periods but not as much in D-block elements because of their (n-1)d sublevel electrons added that protect the outer electrons from the nucleus


the atomic radii of hafnium is slightly less than that of zirconium, but elements in the sixth period following hafnium increase in the usual manner

Ionization Energy

D-Block & F-Block-

Generally increase across the periods, increase down each group

Ion Formation & Ionic Radii

D-Block & F-Block-

Electrons in the highest occupied sublevels are always removed first



all d-block elements electronegativities are between 1.1 and 2.54. Electronegativity values increase as radii decrease


all f-block elements electronegativities are between 1.1 and 1.5