Identify at least two ions that have the following ground-state electron configurations: (a) \([\mathrm{Ar}] ;\) (b) \([\mathrm{Ar}] 3 d^{5}\); (c) \([\mathrm{Kr}] 5 s^{2} 4 d^{10}\)

Using these configurations is particularly useful in predicting and explaining the chemical behavior of elements. Noble gases are nonreactive due to their full outer electron shells, providing a stable configuration that other elements strive to achieve through ion formation. By gaining or losing electrons to emulate noble gas electron configurations, elements attain chemically stable, energetically favored configurations. This is a key factor in the formation of ions.

Understanding the periodic table trends is crucial in predicting the chemical and physical properties of elements, including their tendencies to form ions.

Elements exhibit periodic trends in electronegativity, atomic size, and ionization energy, which influence their ability to gain or lose electrons. For example, moving from left to right across a period, atomic size decreases and electronegativity increases. This means elements towards the right are more likely to gain electrons and form anions, whereas elements on the left tend to lose electrons and form cations.

Going down a group, atoms typically increase in size due to the addition of electron shells. This makes it easier for the outermost electrons to be removed, facilitating the formation of cations. Thus, the position of an element on the periodic table can often help predict whether it will more readily form a cation or an anion, and what the resulting electron configuration might be.