Problem 171 - A Complete Resource Book in Chemistry for JEE Main

Question

Which one of the following constitutes a group of the isoelectronic species? \(\quad\) (a) \(\mathrm{C}_{2}^{2-}, \mathrm{O}_{2}^{-}, \mathrm{CO}, \mathrm{NO}\) (b) \(\mathrm{NO}^{+}, \mathrm{C}^{2}{\underline{\phantom{xx}}}_{2}, \mathrm{C}^{\mathrm{N}-}, \mathrm{N}_{2}\) (c) \(\mathrm{CN}^{-}, \mathrm{N} 2, \mathrm{O}_{2}^{2-}, \mathrm{C}_{2}^{\frac{2}{2}}\) (d) \(\mathrm{N}_{2}, \mathrm{O}_{2}^{-}, \mathrm{NO}^{+}, \mathrm{CO}\)

Step-by-Step Solution

Answer

Option (b) \( \mathrm{NO}^{+}, \mathrm{C}^{2}{\underline{\phantom{xx}}}_{2}, \mathrm{CN}^{-}, \mathrm{N}_{2} \) are isoelectronic species.

Step 1: Understand Isoelectronic Species

Isoelectronic species are atoms, molecules, or ions that have the same number of electrons. Therefore, to determine which group is isoelectronic, we need to count the number of electrons in each species and compare them to each other.

Step 2: Calculate Electrons for Option (a)

- For \( \mathrm{C}_{2}^{2-} \): Each carbon atom has 6 electrons, so 2 carbon atoms have 12 electrons. The \(2-\) charge adds 2 electrons, totaling 14 electrons.- For \( \mathrm{O}_{2}^{-} \): Each oxygen atom has 8 electrons, so 2 oxygens have 16 electrons. The \(1-\) charge adds 1 electron, totaling 17 electrons.- For \( \mathrm{CO} \): Carbon has 6 electrons and oxygen has 8 electrons, totaling 14 electrons.- For \( \mathrm{NO} \): Nitrogen has 7 electrons and oxygen has 8 electrons, totaling 15 electrons.

Step 3: Calculate Electrons for Option (b)

- For \( \mathrm{NO}^{+} \): Nitrogen has 7 electrons and oxygen has 8 electrons. The \(1+\) charge subtracts 1 electron, totaling 14 electrons.- For \( \mathrm{C}^{2}{\underline{\phantom{xx}}}_{2} \): Each carbon atom has 6 electrons, so 2 carbon atoms have 12 electrons.- For \( \mathrm{CN}^{-} \): Carbon has 6 electrons and nitrogen has 7 electrons. The \(1-\) charge adds 1 electron, totaling 14 electrons.- For \( \mathrm{N}_{2} \): Each nitrogen atom has 7 electrons, so 2 nitrogen atoms have 14 electrons.

Step 4: Calculate Electrons for Option (c)

- For \( \mathrm{CN}^{-} \): Carbon has 6 electrons, nitrogen has 7 electrons. The \(1-\) charge adds 1 electron, totaling 14 electrons.- For \( \mathrm{N}_{2} \): Each nitrogen atom has 7 electrons, totaling 14 electrons.- For \( \mathrm{O}_{2}^{2-} \): Each oxygen atom has 8 electrons. The \(2-\) charge adds 2 electrons, totaling 18 electrons.- For \( \mathrm{C}_{2}^{\frac{2}{2}} \): Assuming it means \( \mathrm{C}_{2} \), each carbon atom has 6 electrons, totaling 12 electrons.

Step 5: Calculate Electrons for Option (d)

- For \( \mathrm{N}_{2} \): Each nitrogen atom has 7 electrons, so 2 nitrogen atoms have 14 electrons.- For \( \mathrm{O}_{2}^{-} \): Each oxygen atom has 8 electrons, so 2 oxygens have 16 electrons. The \(1-\) charge adds 1 electron, totaling 17 electrons.- For \( \mathrm{NO}^{+} \): Nitrogen has 7 electrons, oxygen has 8 electrons. The \(1+\) charge subtracts 1 electron, totaling 14 electrons.- For \( \mathrm{CO} \): Carbon has 6 electrons and oxygen has 8 electrons, totaling 14 electrons.

Step 6: Identify the Isoelectronic Group

By comparing the electron counts, option (b) is the only group where every species has 14 electrons: \( \mathrm{NO}^{+} \), \( \mathrm{C}^{2}{\underline{\phantom{xx}}}_{2} \), \( \mathrm{CN}^{-} \), and \( \mathrm{N}_{2} \). Therefore, these species are isoelectronic.

Key Concepts

Electron Configuration

Electron configuration is a way to describe the arrangement of electrons within an atom or molecule. It tells us how electrons are distributed in an atom's orbitals, which are the paths they follow around the nucleus. This distribution plays a vital role in predicting the chemical behavior of elements. For example, for carbon, the electron configuration is \( 1s^2 2s^2 2p^2 \), indicating that it has a total of 6 electrons spread across different energy levels and orbitals.

Knowing an element's electron configuration is essential because it helps chemists determine how atoms in different elements will bond to form molecules. Electrons in the outermost shell, known as valence electrons, are typically the ones involved in forming chemical bonds.

• Electron configurations are used to predict the likelihood of an atom gaining or losing electrons to achieve a stable configuration resembling a noble gas.
• The concept of isoelectronic species arises when different atoms or ions have exactly the same electron configuration.
• This usually results in them exhibiting similar chemical behaviors.

Molecular Ions

Molecular ions are formed when atoms lose or gain electrons, resulting in a charged molecule. Positively charged ions are called cations, which occur when molecules lose electrons, while negatively charged ions are called anions, which occur from gaining electrons.

Understanding molecular ions is essential for determining how chemical reactions occur because they directly affect the forces between molecules. Molecular ions can cause attractions that lead to bonding or reaction processes. For example, in option (b), \( \mathrm{NO}^{+} \) is a molecular ion contributing to the isoelectronic set.

• Molecular ions play a crucial role in the stability of molecules and compound formation.
• They are key to reactions in both organic and inorganic chemistry.
• The presence of ions in a substance can change its properties, such as melting point and solubility.

Chemical Bonding

Chemical bonding is the force that holds atoms together within a molecule. It involves the attraction between the electrons and the nuclei of different atoms. There are several types of chemical bonds, including covalent, ionic, and metallic bonds.

Covalent bonds happen when two atoms share one or more pairs of electrons, providing a stable balance of attractive and repulsive forces between atoms. For instance, in \( \mathrm{CO} \) (carbon monoxide), a covalent bond consists of a shared electron pair between carbon and oxygen. Ionic bonds, on the other hand, occur when an electron is transferred from one atom to another, resulting in positive and negative ions that attract each other.

• Understanding chemical bonding helps explain molecular structures and reactions.
• Isoelectronic species often have similar bonding structures because they have the same number of electrons.
• Variations in bonding determine the physical and chemical properties of the material.