Question

Which one of the following sets of ions represents a collection of isoelectronic species? (a) \(\mathrm{K}^{+}, \mathrm{Cl}^{-}, \mathrm{Ca}^{2+}, \mathrm{Sc}^{3+}\) (b) \(\mathrm{Ba}^{2+}, \mathrm{Sr}^{2+}, \mathrm{K}^{+}, \mathrm{Ca}^{2+}\) (c) \(\mathrm{N}^{3-}, \mathrm{O}^{2}, \mathrm{~F}^{-}, \mathrm{S}^{2-}\) (d) \(\mathrm{Li}^{+}, \mathrm{Na}^{+}, \mathrm{Mg}^{2+}, \mathrm{Ca}^{2+}\)

Short answer

Option (a) is the correct answer as all ions are isoelectronic with 18 electrons.

Step-by-step solution

Define Isoelectronic Species

Isoelectronic species are atoms and ions that have the same number of electrons. To find isoelectronic species among given ions, we need to determine the number of electrons in each ion.

Count Electrons for Option (a)

For option (a), calculate the electrons in each ion: - \(\text{K}^{+}\) has 18 electrons (19 - 1), - \(\text{Cl}^{-}\) has 18 electrons (17 + 1), - \(\text{Ca}^{2+}\) has 18 electrons (20 - 2), - \(\text{Sc}^{3+}\) has 18 electrons (21 - 3). All ions have 18 electrons, therefore they are isoelectronic.

Count Electrons for Option (b)

For option (b), calculate the electrons in each ion: - \(\text{Ba}^{2+}\) has 54 electrons (56 - 2), - \(\text{Sr}^{2+}\) has 36 electrons (38 - 2), - \(\text{K}^{+}\) has 18 electrons (19 - 1), - \(\text{Ca}^{2+}\) has 18 electrons (20 - 2). These ions do not have the same number of electrons.

Count Electrons for Option (c)

For option (c), calculate the electrons in each ion: - \(\text{N}^{3-}\) has 10 electrons (7 + 3), - \(\text{O}^{2-}\) has 10 electrons (8 + 2), - \(\text{F}^{-}\) has 10 electrons (9 + 1), - \(\text{S}^{2-}\) has 18 electrons (16 + 2). Not all ions have the same number of electrons.

Count Electrons for Option (d)

For option (d), calculate the electrons in each ion: - \(\text{Li}^{+}\) has 2 electrons (3 - 1), - \(\text{Na}^{+}\) has 10 electrons (11 - 1), - \(\text{Mg}^{2+}\) has 10 electrons (12 - 2), - \(\text{Ca}^{2+}\) has 18 electrons (20 - 2). These ions do not have the same number of electrons.

Final Answer

Collectively review each option. Only option (a) contains ions that are isoelectronic with 18 electrons each.

Key concepts

Electron Configuration

Electron configuration refers to the way electrons are distributed in an atom or ion. Understanding this concept is crucial for identifying isoelectronic species. Electrons fill orbitals in an atom according to specific rules like the Aufbau principle, Pauli exclusion principle, and Hund's rule. The simplest way to represent electron configurations uses the structure with the notation (e.g., 1s² 2s² 2p⁶). This quickly tells us how many electrons are in each shell and subshell.
Understanding electron configurations gives insight into the chemical properties of an element or ion. For instance, noble gases have complete outer shells, resulting in their stability. When considering ions, the electrons in the outermost shell are often added or removed first. This process modifies the electron configuration, yet the goal is to achieve a stable arrangement, much like the nearest noble gas. Recognizing ions with identical electron configurations helps identify whether they are isoelectronic.

Ion Counting

Ion counting is the process of determining the number of electrons in an ion. This is fundamental when identifying isoelectronic species.
For any neutral atom, the number of electrons equals the atomic number on the periodic table. When an ion is formed, electrons are gained or lost:

• When an ion is negatively charged (an anion), electrons are added.
• When an ion is positively charged (a cation), electrons are subtracted.

For example, a chloride ion ( ext{Cl}^{-}) has 17 protons and one additional electron, making it have 18 electrons in total. In contrast, a potassium ion ( ext{K}^{+}) with an atomic number of 19 loses one electron, resulting in 18 electrons. Ions like these, which have altered their electron counts to become like other stable configurations, are helpful in understanding reactions and interactions in chemistry.

Periodic Table Trends

Periodic table trends involve understanding how elements are organized and how their properties change across groups and periods. Recognizing these trends is beneficial when determining isoelectronic species.
Elements in the same group typically display similar chemical behavior because they have the same number of electrons in their outer shell, making them react similarly. For instance, alkali metals in group 1 have a single electron in their outer shell, which they tend to lose easily, forming cations with a noble gas electron configuration.
As one moves across a period, the number of protons and electrons increases. However, adding electrons to the same shell gradually decreases atomic radius as the nuclear charge pulls electrons closer. This is crucial because smaller atomic radii often mean more tightly held electrons, influencing reactivity. Isoelectronic species may have very different atomic radii and reactivity levels despite having the same number of electrons. These differences often arise from the varying numbers of protons pulling on those electrons.