The size of an atom or an ion is a fundamental property of an element. A consistent pattern observed across all vertical columns, known as groups, is the increase in size as one moves from top to bottom. This systematic expansion holds true for both neutral atoms (atomic radius) and their charged counterparts (ionic radius). This uniform expansion is a direct consequence of adding new layers of electron-containing space, a process that overcomes the stronger positive charge of a growing nucleus.
The Role of Adding Electron Shells
The primary cause for the consistent growth in size is the physical addition of electron shells. As elements are ordered down any group, each subsequent element introduces electrons into a new, higher principal quantum number (\(n\)) level. For instance, elements in the first row fill the first shell (\(n=1\)), and the second row elements fill the second shell (\(n=2\)).
Since each new shell is geometrically farther away from the central nucleus than the last, the overall boundary of the atom expands. The distance of the outermost electrons, known as valence electrons, from the nucleus defines atomic size. Placing these valence electrons into a larger, more distant energy level is the necessary condition for the observed size increase.
Electron Shielding and Effective Nuclear Charge
While adding a new shell provides space for expansion, a mechanism is required to explain why the atom does not contract under the pull of an increasing number of protons. Moving down a group, the atomic number increases, which should lead to a stronger pull on surrounding electrons. This is where the concept of electron shielding becomes important.
The electrons occupying the inner shells, known as core electrons, reside between the positively charged nucleus and the outermost valence electrons. These core electrons repel the valence electrons and partially block the nucleus’s attractive force from reaching the outer layer. This protective effect is termed electron shielding. The net positive attraction experienced by the valence electrons is called the Effective Nuclear Charge (\(Z_{eff}\)).
The \(Z_{eff}\) is calculated by subtracting the shielding constant (S) from the total nuclear charge (Z). As one descends a group, the total nuclear charge (Z) increases significantly, but the number of core electrons (S) increases proportionally because an entire new shell of inner electrons is added. This proportional increase in shielding means that the \(Z_{eff}\) experienced by the valence electrons remains relatively constant, or increases only slightly, down a group. The combination of a new, larger orbital shell and a relatively constant \(Z_{eff}\) allows the atom to expand.
Size Increase in Neutral Atoms
Applying these two principles explains the steady increase in atomic radius down every group for neutral atoms. Comparing lithium (Li) to sodium (Na), the valence electron moves from the second electron shell to the third electron shell. This shift places the outermost electron farther from the nucleus.
The valence electron in sodium experiences a stronger total nuclear charge (11 protons) than the valence electron in lithium (3 protons). However, the increase in core electrons (two in Li to ten in Na) provides a strong shielding effect. This shielding largely neutralizes the increased nuclear attraction, resulting in an effective nuclear charge that is not dramatically different between the two elements.
Since the valence electron resides in a higher principal quantum number shell and is not pulled inward with greater force, the distance between the nucleus and the outermost electron cloud steadily increases. This results in a predictable trend: the atomic radius of neutral atoms grows larger with every step down a group.
Size Increase in Ions
The same fundamental physics governs the trend for ions, resulting in the ionic radius also increasing down a group. When an atom forms an ion, it either gains electrons (anion) or loses electrons (cation). While forming an ion changes its size relative to the parent atom, the addition of a new, higher principal quantum number shell remains the dominant factor controlling size within the group.
When comparing lithium cation (\(\text{Li}^+\)) to sodium cation (\(\text{Na}^+\)), the sodium ion still possesses a greater number of filled electron shells. Specifically, \(\text{Li}^+\) has electrons up to the \(n=1\) shell, while \(\text{Na}^+\) has electrons up to the \(n=2\) shell.
This difference ensures that the ionic radius of \(\text{Na}^+\) is larger than \(\text{Li}^+\), mirroring the trend observed in the neutral atoms. The same logic applies to anions; the fluoride ion (\(\text{F}^-\)) is smaller than the chloride ion (\(\text{Cl}^-\)) because the chloride ion has an extra, larger electron shell. The addition of electron shells down a group uniformly dictates the expansion of the ion’s boundary, maintaining the size increase trend.