Problem 159 - A Complete Resource Book in Chemistry for JEE Main

Question

Which one of the following grouping represents a collection of isoelectronic species? (At. numbers \(\mathrm{Cs}-55, \mathrm{Br}-35)\) (a) \(\mathrm{Na}^{+}, \mathrm{Ca}^{2+}, \mathrm{Mg}^{2+}\) (b) \(\mathrm{N}^{3}, \mathrm{~F}^{-} \mathrm{Na}^{+}\) (c) \(\mathrm{Be}, \mathrm{Al}^{3+}, \mathrm{Cl}^{-}\) (d) \(\mathrm{Ca}^{2+}, \mathrm{Cs}^{+}, \mathrm{Br}\)

Step-by-Step Solution

Answer

Group (b) \(\mathrm{N}^{3-}, \mathrm{F}^{-}, \mathrm{Na}^{+}\) are isoelectronic.

Step 1: Define Isoelectronic Species

Isoelectronic species are atoms, molecules, or ions that have the same number of electrons. Thus, to determine if the given species are isoelectronic, we need to find the number of electrons in each.

Step 2: Determine Electron Count for Each Species in Group (a)

For group (a), \(\mathrm{Na}^{+}\) has 10 electrons because a neutral sodium atom (atomic number 11) loses one electron to become \(\mathrm{Na}^{+}\). \(\mathrm{Ca}^{2+}\) has 18 electrons as a neutral calcium atom (atomic number 20) loses two electrons. \(\mathrm{Mg}^{2+}\) has 10 electrons because a neutral magnesium atom (atomic number 12) loses two electrons. Hence, \(\mathrm{Na}^{+}\) and \(\mathrm{Mg}^{2+}\) are isoelectronic with 10 electrons, but \(\mathrm{Ca}^{2+}\) is not isoelectronic with them.

Step 3: Determine Electron Count for Each Species in Group (b)

In group (b), \(\mathrm{N}^{3-}\) has 10 electrons because a neutral nitrogen atom (atomic number 7) gains three electrons. \(\mathrm{F}^{-}\) has 10 electrons as a neutral fluorine atom (atomic number 9) gains one electron. \(\mathrm{Na}^{+}\) also has 10 electrons. All species in group (b) have the same number of electrons (10), thus they are isoelectronic.

Step 4: Determine Electron Count for Each Species in Group (c)

For group (c), \(\mathrm{Be}\) has 4 electrons because its atomic number is 4. \(\mathrm{Al}^{3+}\) has 10 electrons since a neutral aluminum atom (atomic number 13) loses three electrons. \(\mathrm{Cl}^{-}\) has 18 electrons as a neutral chlorine atom (atomic number 17) gains one electron. These cannot be isoelectronic,

Step 5: Determine Electron Count for Each Species in Group (d)

In group (d), \(\mathrm{Ca}^{2+}\) has 18 electrons. \(\mathrm{Cs}^{+}\) has 54 electrons as a neutral cesium atom (atomic number 55) loses one electron. \(\mathrm{Br}\) has 35 electrons as its atomic number is 35. These are not isoelectronic.

Key Concepts

Electron Configuration

The electron configuration of an atom describes how electrons are distributed in atomic orbitals. For example, in a neutral atom, electrons fill the lowest energy levels or shells available. These configurations not only influence an atom's chemical properties but also determine if an atom becomes isoelectronic with others when it gains or loses electrons.
When atoms form ions, like \(\mathrm{Na}^{+}\), their electron configuration changes to reflect a loss or gain of electrons. Sodium, with atomic number 11, loses one electron to become \(\mathrm{Na}^{+}\), changing its electron configuration from \([\text{Ne}]3s^1\) to \([\text{Ne}]\). Similarly, other ions exhibit changes that determine their chemical behavior and interactions with other species.
Electron configurations play a crucial role in identifying if species are isoelectronic. Isoelectronic species have the same number of electrons, thus sharing a similar electron arrangement. By understanding electron configuration, predicting chemical behavior becomes straightforward.

Ionic Species

Ionic species are atoms that have acquired a positive or negative charge due to the loss or gain of electrons. Unlike neutral atoms, ions are charged entities that play crucial roles in chemical reactions.
Positively charged ions, or cations, result from atoms losing one or more electrons. For instance, \(\mathrm{Na}^{+}\) and \(\mathrm{Ca}^{2+}\) are cations formed from sodium and calcium, respectively. Conversely, negatively charged ions, or anions, like \(\mathrm{F}^{-}\) or \(\mathrm{Cl}^{-}\), occur when atoms gain electrons.
Understanding ionic species is key in investigating isoelectronic collections. By determining how the gain or loss of electrons alters the electron count, we can identify species like \(\mathrm{Na}^{+}\) and \(\mathrm{N}^{3-}\) as having identical number of electrons, making them isoelectronic.

Atomic Number

The atomic number of an element signifies the number of protons in its nucleus. It's a unique identifier that places the element in the periodic table's sequence.
The atomic number not only determines an element's position, but also its inherent properties like electron configuration and chemical reactivity. For instance, sodium (Na) has an atomic number of 11, which informs us that a neutral sodium atom also contains 11 electrons.
When atoms transform into ions, their atomic number remains constant, though their electron count changes due to electron loss or gain. Recognizing atomic numbers helps differentiate similar species and is essential during the identification of isoelectronic groups. Knowing that \(\mathrm{F}^{-}\) and \(\mathrm{N}^{3-}\) have atomic numbers 9 and 7 respectively, alerts us they must gain electrons to become isoelectronic with a neutral neon atom.

Electron Count

Electron count is a term that describes the exact number of electrons present in an atom or ion. This count is vital in the identification of isoelectronic species.
To determine electron count, start with an element’s atomic number, which tells the number of electrons in a neutral atom. Adjust this total for ions by accounting for the electrons lost (cation) or gained (anion).
For example, consider an \(\mathrm{N}^{3-}\) ion. Nitrogen, with an atomic number of 7, gains three electrons to become \(\mathrm{N}^{3-}\), giving it a total electron count of 10. Similar calculations apply to ions like \(\mathrm{Na}^{+}\) and \(\mathrm{F}^{-}\).
By using electron count, identifying isoelectronic species, like those sharing 10 electrons, becomes a straightforward process. This knowledge aids in understanding and predicting the behavior of different ions in chemical systems.