Each of the following elements is capable of forming an ion in chemical reactions. By referring to the periodic table, predict the charge of the most stable ion of each: (a) Be, (b) \(\mathrm{Rb}\), (c) As, \((\mathbf{d})\) In, (e) At.

Consult the periodic table to find the group number for each element. The group number helps us determine the number of valence electrons. (a) Be is in Group 2, so it has 2 valence electrons. (b) Rb is in Group 1, so it has 1 valence electron. (c) As is in Group 15, so it has 5 valence electrons. (d) In is in Group 13, so it has 3 valence electrons. (e) At is in Group 17, so it has 7 valence electrons.

Next, predict the charge of the most stable ion for each element based on the valence electrons. (a) Be tends to lose 2 electrons to have a full outer shell, forming an ion with a \(+2\) charge: \(Be^{2+}\). (b) Rb tends to lose 1 electron to have a full outer shell, forming an ion with a \(+1\) charge: \(Rb^{+}\). (c) As tends to gain 3 electrons to have a full outer shell, forming an ion with a \(-3\) charge: \(As^{3-}\). (d) In tends to lose 3 electrons to have a full outer shell, forming an ion with a \(+3\) charge: \(In^{3+}\). (e) At tends to gain 1 electron to have a full outer shell, forming an ion with a \(-1\) charge: \(At^{-}\). The charges of the most stable ions for the given elements are: (a) \(Be^{2+}\) (b) \(Rb^{+}\) (c) \(As^{3-}\) (d) \(In^{3+}\) (e) \(At^{-}\)