Introduction to the Periodic Table

History and Format of the Periodic Table of the Elements

Dmitri Mendeleev is credited with developing the first periodic table of the elements. His table organized elements by atomic weight. The modern table is organized by atomic number.
Dmitri Mendeleev is credited with developing the first periodic table of the elements. His table organized elements by atomic weight. The modern table is organized by atomic number. Andrey Prokhorov / Getty Images

Dmitri Mendeleev published the first periodic table in 1869. He showed that when the elements were ordered according to atomic weight, a pattern resulted where similar properties for elements recurred periodically. Based on the work of physicist Henry Moseley, the periodic table was reorganized on the basis of increasing atomic number rather than on atomic weight. The revised table could be used to predict the properties of elements that had yet to be discovered.

Many of these predictions were later substantiated through experimentation. This led to the formulation of the periodic law, which states that the chemical properties of the elements are dependent on their atomic numbers.

Organization of the Periodic Table

The periodic table lists elements by atomic number, which is the number of protons in every atom of that element. Atoms of an atomic number may have varying numbers of neutrons (isotopes) and electrons (ions), yet remain the same chemical element.

Elements in the periodic table are arranged in periods (rows) and groups (columns). Each of the seven periods is filled sequentially by atomic number. Groups include elements having the same electron configuration in their outer shell, which results in group elements sharing similar chemical properties.

The electrons in the outer shell are termed valence electrons. Valence electrons determine the properties and chemical reactivity of the element and participate in chemical bonding.

The Roman numerals found above each group specify the usual number of valence electrons.

There are two sets of groups. The group A elements are the representative elements, which have s or p sublevels as their outer orbitals. The group B elements are the nonrepresentative elements, which have partly filled d sublevels (the transition elements) or partly filled f sublevels (the lanthanide series and the actinide series).

The Roman numeral and letter designations give the electron configuration for the valence electrons (e.g., the valence electron configuration of a group VA element will be s2p3 with 5 valence electrons).

Another way to categorize elements is according to whether they behave as metals or nonmetals. Most elements are metals. They are found on the lefthand side of the table. The far right side contains the nonmetals, plus hydrogen displays nonmetal characteristics under ordinary conditions. Elements that have some properties of metals and some of nonmetals are called metalloids or semimetals. These elements are found along a zig-zag line that runs from the upper left of group 13 to the bottom right of group 16. Metals are generally good conductors of heat and electricity, are malleable and ductile, and have a lustrous metallic appearance. In contrast, most nonmetals are poor conductors of heat and electricity, tend to be brittle solids, and can assume any of a number of physical forms. While all of the metals except mercury are solid under ordinary conditions, nonmetals may be solids, liquids, or gases at room temperature and pressure. Elements may be further subdivided into groups. Groups of metals include the alkali metals, alkaline earth metals, transition metals, basic metals, lanthanides, and actinides.

Groups of nonmetals include the nonmetals, halogens, and noble gases.

Periodic Table Trends

The organization of the periodic table leads to recurring properties or periodic table trends. These properties and their trends are:

Ionization Energy - energy needed to remove an electron from a gaseous atom or ion. Ionization energy increases moving left to right and decreases moving down an element group (column).

Electronegativity - how likely an atom is to form a chemical bond. Electronegativity increases moving left to right and decreases moving down a group. The noble gases are an exception, with an electronegativity approaching zero.

Atomic Radius (and Ionic Radius) - a measure of the size of an atom. Atomic and ionic radius decreases moving left to right across a row (period) and increases moving down a group.

Electron Affinity - how readily an atom accepts an electron. Electron affinity increases moving across a period and decreases moving down a group. Electron affinity is nearly zero for noble gases.