The energy of the n = 1 shell also decreases tremendously (the filled 1 s orbital becomes more stable) as the nuclear charge increases. Consequently, the two electrons in the n = 1 shell experience nearly the full nuclear charge, resulting in a strong electrostatic interaction between the electrons and the nucleus. Because the 1 s 2 shell is closest to the nucleus, its electrons are very poorly shielded by electrons in filled shells with larger values of n. The peak for the filled n = 1 shell occurs at successively shorter distances for neon ( Z = 10) and argon ( Z = 18) because, with a greater number of protons, their nuclei are more positively charged than that of helium. ![]() Argon, with filled n = 1, 2, and 3 principal shells, has three peaks. In contrast, neon, with filled n = 1 and 2 principal shells, has two peaks. Because helium has only one filled shell ( n = 1), it shows only a single peak. Each peak in a given plot corresponds to the electron density in a given principal shell. In Ar, the 1 s electrons have a maximum at ≈2 pm, the 2 s and 2 p electrons combine to form a maximum at ≈18 pm, and the 3 s and 3 p electrons combine to form a maximum at ≈70 pm.įigure 3.2.1 also shows that there are distinct peaks in the total electron density at particular distances and that these peaks occur at different distances from the nucleus for each element. In Ne, the 1 s electrons have a maximum at ≈8 pm, and the 2 s and 2 p electrons combine to form another maximum at ≈35 pm (the n = 2 shell). In He, the 1 s electrons have a maximum radial probability at ≈30 pm from the nucleus. This text is adapted from OpenStax Chemistry 2e, Section 6.5: Periodic Variations in Element Properties.\( \newcommand\)įigure 3.2.1 Plots of Radial Probability as a Function of Distance from the Nucleus for He, Ne, and Ar. The greater the nuclear charge, the smaller the radius in a series of isoelectronic ions and atoms. For atoms or ions that are isoelectronic, the number of protons determines the size. ![]() Another isoelectronic series is P 3–, S 2–, Cl –, Ar, K +, Ca 2+, and Sc 3+ (3 s 23 p 6). Examples of isoelectronic species are N 3–, O 2–, F –, Ne, Na +, Mg 2+, and Al 3+ (1 s 22 s 22 p 6). For consecutive elements proceeding down any group, anions have larger principal quantum numbers and, thus, larger radii.Ītoms and ions that have the same electron configuration are said to be isoelectronic. For example, a sulfur atom (3 s 23 p 4) has a covalent radius of 104 pm, whereas the ionic radius of the sulfide anion (3 s 23 p 6) is 170 pm. Both effects (the increased number of electrons and the decreased Z eff) cause the radius of an anion to be larger than that of the parent atom. This results in a greater repulsion among the electrons and a decrease in Z eff per electron. Proceeding down the groups of the periodic table, cations of successive elements with the same charge generally have larger radii, corresponding to an increase in the principal quantum number, n.Īn anion (negative ion) is formed by the addition of one or more electrons to the valence shell of an atom. As electrons are removed from the outer valence shell, the remaining core electrons occupying smaller shells experience a greater effective nuclear charge Z eff and are drawn even closer to the nucleus.Ĭations with larger charges are smaller than cations with smaller charges (e.g., V 2+ has an ionic radius of 79 pm, while that of V 3+ is 64 pm). For example, the covalent radius of an aluminum atom (1 s 22 s 22 p 63 s 23 p 1) is 118 pm, whereas the ionic radius of an Al 3+ (1 s 22 s 22 p 6) is 68 pm. A cation always has fewer electrons and the same number of protons as the parent atom it is smaller than the atom from which it is derived. ![]() Ionic radius is the measure used to describe the size of an ion.
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