The periodic table organizes elements in a grid based on their atomic structure and recurring chemical properties. This arrangement uses two primary classifications: vertical columns and horizontal rows. The horizontal row, known formally as a period, dictates several physical and chemical characteristics of the elements it contains.
The Definition of a Period
A period is one of the seven horizontal rows on the periodic table. The period number indicates the number of principal energy levels, or electron shells, that an atom of that element possesses in its ground state. Elements in the first period have one electron shell, while elements in the seventh period have seven occupied shells.
All elements within the same period share this common number of electron shells. Moving from left to right, each successive element gains one more proton and one more electron. These additional electrons are added to the same outermost energy shell, which keeps the row number constant across the period.
Distinguishing Periods from Groups
The organization of the periodic table relies on the distinction between periods and the vertical columns, which are called groups. Elements in the same period share the same number of electron shells, but they often have widely different chemical behaviors. For example, a period might start with a highly reactive metal and end with an inert noble gas.
In contrast, elements within the same group share the same number of valence electrons (outermost shell electrons). This shared number of valence electrons gives elements in the same group similar chemical properties. While the period defines the atom’s structure through its number of shells, the group defines its chemical behavior through its valence electrons.
Changing Properties Across a Period
Moving from left to right across any horizontal row reveals systematic changes in the elements’ properties, known as periodic trends. As the atomic number increases, the nucleus gains more protons, which increases the positive charge felt by the outermost electrons. This stronger attraction, termed the effective nuclear charge, pulls the electron cloud closer to the nucleus.
This increased pull causes the atomic radius, or size of the atom, to decrease. Similarly, the energy required to remove an electron from the atom, known as ionization energy, increases because the electrons are held more tightly. The electronegativity, which is an atom’s ability to attract electrons in a chemical bond, also increases. These changes reflect a transition from elements that readily lose electrons (metals) on the left to elements that tend to gain electrons (non-metals) on the right.