An understanding of periodic trends is necessary when analyzing and predicting molecular properties and interactions. Common periodic trends include those in ionization energy, atomic radius, and electron affinity. One such trend is closely linked to atomic radii -- ionic radii. Neutral atoms tend to increase in size down a group and decrease across a period. When a neutral atom gains or loses an electron, creating an anion or cation, the atom"s radius increases or decreases, respectively. This module explains how this occurs and how this trend differs from that of atomic radii.

Shielding and Penetration

Electromagnetic interactions between electrons in an atom modify the effective nuclear charge ((Z_eff)) on each electron. Penetration refers to the presence of an electron inside the shell of an inner electron, and shielding is the process by which an inner electron masks an outer electron from the full attractive force of the nucleus, decreasing (Z_eff). Differences in orbital characteristics dictate differences in shielding and penetration. Within the same energy level (indicated by the principle quantum number, n), due to their relative proximity to the nucleus, s-orbital electrons both penetrate and shield more effectively than p-orbital electrons, and p electrons penetrate and shield more effectively than d-orbital electrons. Shielding and penetration along with the effective nuclear charge determine the size of an ion. An overly-simplistic but useful conceptualization of effective nuclear charge is given by the following equation:


(Z) is the number of protons in the nucleus of an atom or ion (the atomic number), and (S) is the number of core electrons.

Figure (PageIndex1) illustrates how this equation can be used to estimate the effective nuclear charge of sodium:

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The Periodic Trend

Due to each atom’s unique ability to lose or gain an electron, periodic trends in ionic radii are not as ubiquitous as trends in atomic radii across the periodic table. Therefore, trends must be isolated to specific groups and considered for either cations or anions.

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Consider the s- and d-block elements. All metals can lose electrons and form cations. The alkali and alkali earth metals (groups 1 and 2) form cations which increase in size down each group; atomic radii behave the same way. Beginning in the d-block of the periodic table, the ionic radii of the cations do not significantly change across a period. However, the ionic radii do slightly decrease until group 12, after which the trend continues (Shannon 1976). It is important to note that metals, not including groups 1 and 2, can have different ionic states, or oxidation states, (e.g. Fe2+ or Fe3+ for iron) so caution must be employed when generalizing about trends in ionic radii across the periodic table.

All non-metals (except for the noble gases which do not form ions) form anions which become larger down a group. For non-metals, a subtle trend of decreasing ionic radii is found across a pegroup theoryriod (Shannon 1976).

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Anions are almost always larger than cations, although there are some exceptions (i.e. fluorides of some alkali metals).