Atomic size, or atomic radius, is a fundamental property that dictates how an element interacts chemically. Precise measurement is challenging because the electron cloud lacks a fixed boundary. Atomic size is typically defined using operational measurements. For instance, the covalent radius is half the distance between the nuclei of two identical atoms joined by a chemical bond. For metals, the metallic radius is used, defined as half the distance between adjacent nuclei in a solid lattice. The periodic table organizes the systematic change in atomic size across all elements.
The Trend Down a Group: Adding Electron Shells
Moving down a vertical column, or group, on the periodic table results in a consistent increase in atomic size. This expansion occurs primarily because each subsequent element adds a new principal quantum shell of electrons. These added shells are progressively further away from the nucleus. The distance between the nucleus and the outermost electrons, which defines the atomic radius, grows substantially with each new shell.
The effect of electron shielding also plays a significant role in the larger size. Inner-shell electrons repel the valence electrons in the outermost shell. This repulsive action effectively blocks the full attractive force of the nucleus from reaching the outer electrons. As more inner shells are added down a group, the screening effect becomes more pronounced.
The combination of a greater number of electron shells and enhanced electron shielding allows the outermost electron cloud to expand. Consequently, valence electrons are not pulled toward the nucleus as strongly. This mechanism explains why Cesium (Cs) at the bottom of Group 1 is significantly larger than Lithium (Li) at the top. The increase in the principal quantum number is the dominant factor driving the increase in atomic size down any group.
The Trend Across a Period: Effective Nuclear Charge
In contrast to the vertical trend, atomic size generally decreases as one moves horizontally from left to right across a period. This pattern occurs even though each element gains both a proton and an electron compared to the one preceding it. The phenomenon is explained by the concept of effective nuclear charge (\(Z_{eff}\)), which is the net positive charge experienced by the outermost valence electrons.
As elements progress across a period, new electrons are added to the same outermost shell. Simultaneously, the number of protons in the nucleus increases, raising the total positive nuclear charge. The inner electrons provide a relatively constant amount of shielding because the number of core shells remains unchanged. Consequently, the valence electrons are poorly shielded from the increasing pull of the nucleus.
The higher nuclear charge, combined with ineffective shielding, results in a steadily increasing \(Z_{eff}\). This stronger net positive charge exerts a greater attractive force on the valence electron cloud. The increased pull draws the outermost electrons closer to the nucleus, causing the atomic radius to shrink. This mechanism explains why elements on the right side of a period, such as the halogens, are considerably smaller than the alkali metals on the far left.
Related Size Comparison: Atoms Versus Ions
The size of an atom changes predictably when it forms an ion. When a neutral atom loses electrons to form a positively charged cation, its size decreases significantly compared to its parent atom. The removal of electrons reduces repulsion among the remaining electrons. Since the same nuclear charge is distributed over fewer electrons, the nucleus exerts a stronger pull, drawing the electron cloud inward and making the cation smaller.
Conversely, when a neutral atom gains electrons to form a negatively charged anion, its size increases. Adding electrons increases the electron-electron repulsion within the outermost shell, forcing the electron cloud to spread out and expanding the atomic radius. Furthermore, the constant nuclear charge is now attracting a larger number of electrons, which decreases the effective pull on any single outer electron. Cations are always smaller than their parent atoms, and anions are always larger.