Problem 124 - A Complete Resource Book in Chemistry for JEE Main

Question

$$ \begin{aligned} &\text { Match the following }\\\ &\begin{array}{ll} \hline \text { Column-I } & \text { Column-II } \\ \hline \begin{array}{ll} \text { (a) } \mathrm{Na}^{+}<\mathrm{F}^{-}<\mathrm{O}^{2-}<\mathrm{N}^{3-} & \text { (p) Electronegativity } \\ \text { (b) } \mathrm{Li}^{+}<\mathrm{Na}^{+}<\mathrm{K}^{+}<\mathrm{Rb}^{+}<\mathrm{Cs}^{+} & \text {(q) } \text { Mobility of } \\ &\text { hydrated ions } \end{array} \\ \begin{array}{ll} \text { (c) } \mathrm{O}<\mathrm{S}<\mathrm{F}<\mathrm{Cl} & \text { (r) Ionic size } \\ \text { (d) } \mathrm{Cl}^{-}<\mathrm{K}^{+}<\mathrm{Ca}^{2+}<\mathrm{Sc}^{3+} & \text { (s) Electron affinity } \\ \hline \end{array} \end{array} \end{aligned} $$

Step-by-Step Solution

Answer

(a) - (r), (b) - (q), (c) - (p), (d) - (r)

Step 1: Analyze Column-I (a)

The order given is \( \mathrm{Na}^{+}<\mathrm{F}^{-}<\mathrm{O}^{2-}<\mathrm{N}^{3-} \). This represents the increasing ionic size. As the number of electrons increases in the same isoelectronic series, the effective nuclear charge decreases, leading to an increase in ionic radius. Thus, Column I (a) matches Column II (r) - Ionic Size.

Step 2: Analyze Column-I (b)

The order \( \mathrm{Li}^{+}<\mathrm{Na}^{+}<\mathrm{K}^{+}<\mathrm{Rb}^{+}<\mathrm{Cs}^{+} \) is according to increasing atomic number in the alkali metals. These represent the hydrated ions, as the smaller the ion, the larger the hydration energy, which inversely affects mobility. Hence, Column I (b) matches Column II (q) - Mobility of hydrated ions.

Step 3: Analyze Column-I (c)

The order \( \mathrm{O}<\mathrm{S}<\mathrm{F}<\mathrm{Cl} \) suggests increasing electronegativity. Fluorine is the most electronegative element, followed by chlorine among these elements. Thus, Column I (c) corresponds to Column II (p) - Electronegativity.

Step 4: Analyze Column-I (d)

The given order \( \mathrm{Cl}^{-}<\mathrm{K}^{+}<\mathrm{Ca}^{2+}<\mathrm{Sc}^{3+} \) shows decreasing ionic radii, as these ions have the same electron configuration but an increasing number of protons, thus a stronger effective nuclear charge and decreasing size. This matches with Column II (r) - Ionic Size.

Key Concepts

Ionic Size

Ionic size refers to the size of an ion, which is determined by the radius of the ion. The radius can vary due to several factors, such as the number of electrons and the effective nuclear charge. In an isoelectronic series, ions with more electrons typically have a larger radius. This is because the effective nuclear charge—the net positive charge experienced by an electron in the atom—decreases as more electrons are added. For example, in the series \( \text{Na}^+ < \text{F}^- < \text{O}^{2-} < \text{N}^{3-} \), each ion has the same number of electrons, but \( \text{Na}^+ \) has the smallest ionic size. This is due to the higher effective nuclear charge within its nucleus, pulling the electrons closer and resulting in a smaller radius. Conversely, \( \text{N}^{3-} \) has the smallest nuclear charge, meaning the electrons are less tightly held, and its ionic size is the largest of the group. Understanding these trends is crucial in fields like chemistry and materials science, where ionic interactions often dictate the properties of compounds.

Electronegativity

Electronegativity is a measure of an atom's ability to attract and hold onto electrons within a chemical bond. This property plays a critical role in determining how atoms interact in molecules. Generally, electronegativity increases across a period from left to right on the periodic table and decreases down a group.Fluorine is known as the most electronegative element due to its position on the periodic table, followed closely by oxygen and chlorine. In the list \( \text{O} < \text{S} < \text{F} < \text{Cl} \), the trend of increasing electronegativity is evident. Chlorine, being one row down from fluorine, does not top fluorine in terms of electronegativity, but it ranks higher compared to oxygen and sulfur, due to its stronger pull on shared electron pairs relative to its atomic radius and effective nuclear charge. Electronegativity differences are key in predicting bond types; for example, a large difference typically results in ionic bonds, whereas smaller differences are associated with covalent bonds. This concept is fundamental in predicting the chemical behavior of atoms.

Mobility of Hydrated Ions

The mobility of hydrated ions is a concept rooted in how ions move through a solution. Ions are surrounded by a layer of water molecules, which is known as the hydration shell. The size of this shell influences the ion's mobility: smaller ions tend to attract more water molecules, forming a larger hydration shell, which can hinder their movement through a solution. In a series such as \( \text{Li}^+ < \text{Na}^+ < \text{K}^+ < \text{Rb}^+ < \text{Cs}^+ \), lithium ions have the most pronounced hydration shell due to their small size, resulting in reduced mobility compared to larger ions like cesium. This is because the hydration energy, or the energy released when ions are surrounded by water molecules, is highest for smaller ions, making their hydrated shells thicker and movement through the solvent more difficult. This principle is particularly important in electrolytes and affects how energy is conducted in biological systems and batteries.

Electron Affinity

Electron affinity is the change in energy when an electron is added to a neutral atom in the gaseous state to form a negative ion. This property varies across the periodic table and generally increases across a period and decreases down a group. Atoms that exhibit high electron affinity values are more likely to accept electrons, forming stable anions. This is why non-metals, which are found on the right side of the periodic table, typically have higher electron affinity values than metals. Moreover, the electron affinity is influenced by atomic size and nuclear charge. Smaller atoms with a stronger effective nuclear charge welcome additional electrons more readily. Chlorine, for example, has a higher electron affinity than many other elements because it can efficiently accommodate an extra electron, completing its valence shell and stabilizing its electronic structure. Understanding electron affinity helps in grasping how elements interact at a molecular level and is a fundamental concept in predicting reactivity and stability of ions in chemical reactions.