Problem 67 - A Complete Resource Book in Chemistry for JEE Main

Question

The correct sequence of the ionic radii of the following is (a) \(\mathrm{I}>\mathrm{S}^{2-}>\mathrm{Cl}^{-}>\mathrm{O}^{2-}>\mathrm{F}^{-}\) (b) \(\mathrm{S}^{2-}>\mathrm{I}^{->}>\mathrm{O}^{2-}>\mathrm{Cl}>\mathrm{F}^{-}\) (c) \(\mathrm{I}>\mathrm{Cl}>\mathrm{S}^{2-}>\mathrm{O}^{2->} \mathrm{F}^{-}\) (d) \(\mathrm{I}^{-}>\mathrm{S}^{2}>\mathrm{Cl}^{-}>\mathrm{F}^{-}>\mathrm{O}^{2-}\)

Step-by-Step Solution

Answer

The correct sequence is option (d): I⁻ > S²⁻ > Cl⁻ > F⁻ > O²⁻.

Step 1: Understanding Ionic Radii

Ionic radii refer to the size of an ion in a crystal lattice. Anions (negative ions) are generally larger than their neutral atoms because they gain electrons, increasing electron-electron repulsion and leading to a larger size.

Step 2: Analyze Given Options

We have been provided with four sequences: (a), (b), (c), and (d). Our task is to determine which sequence is correctly ordered from largest to smallest based on the ionic radii.

Step 3: Compare Ionic Radii

Typically, within the same group (column) of the periodic table, as the atomic number increases, the size of the ion increases. Thus, for halogens: Iodide (I⁻) is larger than Chloride (Cl⁻) and Fluoride (F⁻). Similarly, within the same group of chalcogens, Sulfide (S²⁻) is larger than Oxide (O²⁻).

Step 4: Evaluate Each Option

- Option (a): I > S²⁻ > Cl⁻ > O²⁻ > F⁻ - Option (b): S²⁻ > I⁻ > O²⁻ > Cl > F⁻ - Option (c): I > Cl > S²⁻ > O²⁻ > F⁻ - Option (d): I⁻ > S²⁻ > Cl⁻ > F⁻ > O²⁻ Now, compare each sequence with expected ionic sizes. The general size order based on ion charges and periodic trends is I⁻ > S²⁻ > Cl⁻ > O²⁻ > F⁻.

Step 5: Identify Correct Sequence

Option (d): I⁻ > S²⁻ > Cl⁻ > F⁻ > O²⁻ follows the established trends in ionic radii. Iodide, being a halogen with the largest ionic size in Group 17, is followed by the large Sulfide ion, and then Chloride. Fluoride is smaller than Chloride, and Oxide is smallest among these due due to its high negative charge density and smallest periodic position.

Key Concepts

Ionic Size Trends

Ionic size trends refer to the changes in the size of ions as you move across the periodic table or down a group. When analyzing ionic sizes, keep in mind that cations (positively charged ions) are smaller than their neutral counterparts. This results from loss of electrons, reducing electron-electron repulsion, and leading to a smaller size. Conversely, anions (negatively charged ions), like those discussed in the exercise, grow larger than their neutral atoms due to the gain of electrons.

• Increased electron density results in enhanced electron-electron repulsion.
• This repulsion causes the electron cloud to expand, enlarging the anion.

Understanding these fundamentals helps to predict the sizes of ions relative to their positions in the periodic table.

Periodic Table Trends

Periodic table trends are valuable tools for predicting the properties of elements, including ionic radii. Elements in the same group, or vertical column of the periodic table, share similar properties but often exhibit predictable changes in size.

• As you move down a group, ionic size generally increases. This is because additional electron shells are added, increasing atomic radius.
• When moving across a period from left to right, ionic size tends to decrease as electrons are added to the same shell, causing the effective nuclear charge to increase and draw electrons closer to the nucleus.

These variations explain the differences in size among ions. For example, Iodide (I⁻) ions are larger than Chloride (Cl⁻) ions because they reside further down the group on the periodic table.

Anions vs Neutral Atoms

The contrast between anions and their neutral atoms is key in understanding why ions have different sizes. Anions are formed when atoms gain additional electrons, leading to an increase in size due to augmented electron-electron interaction.

• In a neutral atom, there's a balance between protons and electrons, allowing a compact size.
• In anions, the extra electrons disrupt this balance, increasing repulsion forces within the electron cloud.
• This results in a larger atomic radius compared to the neutral atom.

A clear instance is the comparison of a Sulfur atom to a Sulfide ion (S²⁻). The sulfide ion is notably larger due to extra electrons increasing electron cloud repulsion. Understanding this distinction sheds light on why some ions possess larger radii than their atomic forms.