Question

Match the column I with column II and mark the appropriate choice. $$ \begin{array}{|l|l|l|l|} \hline {\begin{array}{c} \text { Column I } \\ \text { (Atom) } \end{array}} & &{\begin{array}{c} \text { Column II } \\ \text { (No. of unpaired electrons) } \end{array}} \\ \hline \text { (A) } & { }_{15} \mathrm{P} & \text { (i) } & \text { 6 unpaired electrons } \\ \hline \text { (B) } & { }_{24} \mathrm{Cr} & \text { (ii) } & \text { 2 unpaired electrons } \\ \hline \text { (C) } & { }_{26} \mathrm{Fe} & \text { (iii) } & \text { 3 unpaired electrons } \\ \hline \text { (D) } & { }_{14} \mathrm{Si} & \text { (iv) } & \text { 4 unpaired electrons } \\ \hline \end{array} $$ (a) (A) \(\rightarrow\) (ii), (B) \(\rightarrow(\mathrm{i}),(\mathrm{C}) \rightarrow\) (iii), (D) \(\rightarrow\) (iv) (b) (A) \(\rightarrow\) (i), (B) \(\rightarrow\) (iii), (C) \(\rightarrow\) (ii), (D) \(\rightarrow\) (iv) (c) (A) \(\rightarrow\) (iii), (B) \(\rightarrow(\mathrm{i}),(\mathrm{C}) \rightarrow\) (iv), (D) \(\rightarrow\) (ii) (d) (A) \(\rightarrow\) (iv), (B) \(\rightarrow\) (ii), (C) \(\rightarrow\) (i), (D) \(\rightarrow\) (iii)

Short answer

(A) \(\rightarrow\) (iii), (B) \(\rightarrow\) (i), (C) \(\rightarrow\) (iv), (D) \(\rightarrow\) (ii)

Step-by-step solution

Calculate Unpaired Electrons for Phosphorus (P)

For element P with atomic number 15, its electron configuration is 1s^2 2s^2 2p^6 3s^2 3p^3. Looking at the highest energy level (3p), we have 3 unpaired electrons.

Calculate Unpaired Electrons for Chromium (Cr)

For element Cr with atomic number 24, the electron configuration is an exception and written as [Ar] 3d^5 4s^1 instead of [Ar] 3d^4 4s^2 due to extra stability provided by half-filled d-orbital. Therefore, there are 6 unpaired electrons in total.

Calculate Unpaired Electrons for Iron (Fe)

For element Fe with atomic number 26, the electron configuration is [Ar] 3d^6 4s^2. When the 3d subshell is considered, there are 4 unpaired electrons.

Calculate Unpaired Electrons for Silicon (Si)

Element Si with atomic number 14 has the electron configuration 1s^2 2s^2 2p^6 3s^2 3p^2. In the 3p orbital, there are 2 unpaired electrons.

Match Columns I and II

Match each element to the corresponding number of unpaired electrons: \( _{15}\mathrm{P} \) has 3 unpaired electrons, \( _{24}\mathrm{Cr} \) has 6 unpaired electrons, \( _{26}\mathrm{Fe} \) has 4 unpaired electrons, and \( _{14}\mathrm{Si} \) has 2 unpaired electrons. Comparing with options given, the matches are (A) \(\rightarrow\) (iii), (B) \(\rightarrow\) (i), (C) \(\rightarrow\) (iv), (D) \(\rightarrow\) (ii).

Key concepts

Electron Configuration

Understanding electron configuration is crucial for predicting the chemical behavior of atoms. It's the way electrons are distributed in atomic orbitals. Electrons fill orbitals starting with the lowest energy level, following the Aufbau principle, Hund's rule, and the Pauli exclusion principle.

Let's consider the electron configuration of Phosphorus (P), which has an atomic number of 15. The configuration follows the pattern of 1s², 2s², up to 3p³, depicting electrons filling in orbitals from lower to higher energy levels. Likewise, transition metals like Chromium (Cr) can have unusual configurations due to the relative energy levels of the 3d and 4s orbitals. For instance, Cr is [Ar] 3d⁵ 4s¹, not the expected [Ar] 3d⁴ 4s², because half-filled d orbitals confer extra stability.

These configurations directly influence the magnetic properties of an element, as the number of unpaired electrons determines its magnetic nature. Elements with unpaired electrons are paramagnetic, while those fully paired are diamagnetic. Thus, knowing an element's electron configuration allows us to predict its reactivity and other chemical properties.

Atomic Structure

The atomic structure encompasses protons, neutrons, and electrons within an atom, defining its identity and properties. At the center is the nucleus, containing protons and neutrons, around which electrons orbit in specific regions called orbitals.

This structure is intricately linked to electron configuration. Orbitals are filled according to known principles and result in layers or shells of electrons. For example, Silicon (Si), atomic number 14, has its electrons arranged as 1s² 2s² 2p⁶ 3s² 3p². The outermost orbitals, in this case the 3p orbital, hold the valence electrons and have the most impact on the atom's chemical behavior. The atomic structure not only dictates the distribution of these electrons but also their interaction with other atoms, influencing the bond formation and overall chemistry of the element.

D-block Elements

D-block elements, also known as transition metals, are found in groups 3-12 of the periodic table and are characterized by their partially filled d orbitals. These elements exhibit unique electron configurations and a wide range of oxidation states, leading to complex chemical behavior.

For example, Iron (Fe) is a d-block element with an atomic number of 26, and its electron configuration is [Ar] 3d⁶ 4s². The presence of unpaired electrons in its d orbitals makes iron paramagnetic. Upon interaction with a magnetic field, unpaired electrons align with the field, making the atom behave like a small magnet.

Transition metals like chromium and iron demonstrate the concept of variable valency and the formation of colored compounds, which is not commonly observed in s or p block elements. These unique properties arise from the complex atomic and electronic structures of d-block elements, making them an integral part of inorganic chemistry and various applications, including catalysis and material science.