Stable Electron Configurations and Charges on Ions

Polar water molecules are strongly attracted to each other. 

I AIMS To learn about stable electron configurations. To iearn to predict the formuias of ionic compounds. 

We have seen many times that when a metal and a nonmetal react to form an ionic compound, the metal atom loses one or more electrons to the non-metal. In Chapter 5, where binary ionic compounds were introduced, we saw that in these reactions. Group 1 metals always form 1 + cations. Group 2 metals always form 2+ cations, and aluminum in Group 3 always forms a 3+ cation. For the nonmetals, the Group 7 elements always form 1-anions, and the Group 6 elements always form 2- anions. This is further illustrated in Table 11.2. 

Atoms in stable compounds almost always have a noble gas electron configuration. 

Representative (main-group) metals form ions by losing enough electrons to achieve the configuration of the previous noble gas (that is, the noble gas that occurs before the metal in question on the periodic table). For example, note from the periodic table inside the front cover of the text that neon is the noble gas previous to sodium and magnesium. Similarly, helium is the noble gas previous to lithium and beryllium. 

OBJECTIVES To learn about Stable electron configurations. To learn to predict the formulas of ionic compounds. 

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We have seen many times that when a metal and a nonmetal react to form an ionic compound, the metal atom loses one or more electrons to the nonmetal. 

Notice something very interesting about the ions in Table 12.2 all have the electron configuration of neon, a noble gas. That is, sodium loses its one valence electron (the 3s) to form Na, which has a [Ne] electron configuration. Likewise, Mg loses its two valence electrons to form Mg, which also has a [Ne] electron configuration. On the other hand, the nonmetal 

Representative (main-group) metals form ions by losing enough electrons to achieve the configuration of the previous noble gas (that is, the noble gas that occurs before the metal in question on the periodic table). 

For example, note from the periodic table inside the back cover of this book that neon is the noble gas before sodium and magnesium. 

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Metals can lose these outer electrons and form positive ions the charge is due to the higher number of protons (+) that remain relative to the number of electrons. The number of electrons lost is equal to the group number in the periodic table and the resultant ion has a stable electronic configuration like that of an inert gas. The net charge on the ion is written to the right of the chemical symbol as a small superscript. 

The metals in Group IIA have two valence electrons, both in an s-subshell outside the stable noble gas core. It should not be surprising that each forms an ion with a 2+ charge, the result of losing both valence electrons. Be2+, Mg2+, and Ca2+ each have the stable electronic configuration of a noble gas. The l[Kr]3d °) configuration of the strontium ion, Srz+, with a completely tilled 3d-subshell outside the [Kr] noble gas core, is called a pseudo-noble gas core. Pseudo-noble gas cores are very stable, too. Don t be confused by this. Focus on the valence electrons, those in the highest occupied principal shell, the n = 5 shell in strontium. They are the chemically important electrons. 

This is larger than the corresponding value for Li (57 kJ mol" ) but substantially smaller than the value for F (333kJmol" ). The hydride ion H" has the same electron configuration as helium but is much less stable because the single positive charge on the proton must now control the 2 electrons. The hydride ion is thus readily deformable and this constitutes a characteristic feature of its structural chemistry (see p. 66). 

The charges on the chlorine, potassium, and calcium ions result from a strong tendency of valence electrons to adopt the stable configuration of the inert gases, with completely filled electronic shells. Notice that the 3 ions have electronic configurations identical to that of inert argon. 

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