Chapter 5: The Periodic Law

Section 3: Electron Configuration and Periodic Properties


Atomic Radius: may be defined as one-half the distance between the nuclei of identical atoms that are bonded together
Ion: an atom, radical, or molecule that has gained or lost one or more electrons and has a negative or positive charge
Ionization: the process of adding or removing electrons from an atom or molecule, which gives the atom or molecule a net charge
Ionization Energy: the energy required to remove an electron from an atom or ion
Electron Affinity: the energy needed to remove an electron from a negative ion to form a neutral atom or molecule
Cation: an ion that has a positive charge
Anion: an ion that has a negative charge
Valence Electrons: an electron that is found in the outermost shell of an atom and that determines the atom's chemical properties
Electronegativity: a measure of the ability of an atom in a chemical compound to attract electrons

Atomic Radii

* In order to estimate the size of an atom, the conditions under which the same atom exists must be specified.

* The trend to smaller atoms across a period is caused by the increasing positive charge of the nucleus

* In general, the atomic radii of the main-group elements increase down a group.

Ionization Energy

* An electron can be removed from an atom if enough energy is supplied.

* In general, ionization energies of the main-group elements increase across each period.

* Among the main-group elements, ionization energies generally decrease down the group.

* With sufficient energy, electrons can be removed from positive ions as well as neutral atoms.

* The energies for removal of additional electrons from an atom are referred to as The second ionization energy (IE2), third ionization energy (IE3), and so on.

* Each successive electron removed from an ion feels an increasingly stronger effective nuclear charge (the nuclear charge minus the electron shielding).

Ionic Radii

* Metals at the left tend to form cations and the nonmetals at the upper right tend to form anions.
* Cationic radii decrease across a period because the electron cloud shrinks due to the increasing nuclear charge acting on the electrons in the same main energy level.
* Starting with Group 15, in which atoms assume noble-gas configurations by gaining three electrons.
* Anionic radii decrease across each period for the elements in Groups 15-18
* Just as there is a gradual increase of atomic radii down a group, there is also a gradual increase of ionic radii.

Valence Electrons

* Chemical compounds form because electrons are lost, gained, or shared between atoms.
* The electrons in the highest energy level are most subject to the influence of nearby atoms or ions.
* Valence electrons are often located in incompletely filled main-energy levels.
* The inner electrons are in filled energy levels and are held too tightly by the nucleus to be involved in compound formation.
* Groups 13-18 have a number of valence electrons equal to the group number minus 10.


* In many compounds, the negative charge of the valence electrons in concentrated closer to one atom than to another.
* Linus Pauling devised a scale of numerical values reflecting the tendency of an atom to attract electrons.
* Electronegativities tend to increase across each period, although there are expectations.
* Electronegativities tend to either decrease down a group or remain about the same.

Electron Affinity

Electron affinity is the opposite of electronegativity. As EN increases EA decreases and visa versa.