The periodic table organizes the chemical elements based on the periodic law, which shows that properties recur when elements are ordered by increasing atomic number. The table is structured with horizontal rows called periods and vertical columns known as groups. Elements within a period do not share the same chemical properties as those in a group, but they possess a fundamental structural similarity that governs how their characteristics change across the row.
Shared Feature The Number of Electron Shells
The defining commonality for all elements in the same period is the number of principal energy levels, or electron shells, that contain electrons. These energy levels represent the general location of an atom’s electrons around its nucleus. For any given period, the period number directly corresponds to the total number of occupied electron shells the atoms possess. For instance, elements in Period 1, such as Hydrogen and Helium, occupy only the first shell, while all elements in Period 4 have electrons in four distinct principal energy levels.
As one moves from left to right across any period, electrons are consistently added to the same outermost electron shell. For example, in Period 3, both Sodium (Na) and Chlorine (Cl) have their electrons spread across three principal energy shells. The key difference is that each step to the right adds one more proton to the nucleus and one more electron to this same outer shell. This constancy dictates the overall size and the way the nucleus exerts its influence over the electrons.
Resulting Trends Changes in Atomic Size and Energy
The constant number of electron shells, combined with the increasing number of protons, creates predictable changes in physical properties across a period. Moving from left to right, the atomic number increases, meaning the nucleus gains a stronger positive charge due to the addition of protons. This increased positive charge, known as the effective nuclear charge, exerts a greater attractive force on the electrons in the constant number of shells. This stronger pull causes the overall size of the atom, or its atomic radius, to decrease across the period.
This tightening of the electron cloud impacts the energy required to remove an electron from the atom, a property called ionization energy. Because the outer electrons are held more tightly by the increasing nuclear charge, it becomes progressively harder to pull one away. Consequently, the ionization energy increases when moving from the left side of the period to the right side. Electronegativity is the measure of an atom’s tendency to attract a bonding pair of electrons toward itself. This property also increases across a period because the stronger nuclear pull makes the atom more effective at drawing in external electrons.
The Shift in Chemical Character
The changes in atomic size, ionization energy, and electronegativity translate directly into a shift in the chemical character of the elements across a period. Elements on the far left, such as the alkali metals, have low ionization energy and readily lose their single outermost electron. This tendency makes them highly reactive metals that form positive ions and participate in ionic bonding.
Transition and Non-Metals
As one progresses across the period, the elements transition through metalloids, which exhibit properties of both metals and non-metals, such as Silicon in Period 3. Elements toward the right side, like the halogens, have high electronegativity and a strong tendency to gain electrons to complete their outer shell. These highly reactive non-metals often form negative ions or engage in covalent bonding by sharing electrons.
Noble Gases
The period culminates with the noble gases on the far right. They have a full outer shell, resulting in high stability and minimal chemical reactivity.