Problem 42 - A Complete Resource Book in Chemistry for JEE Main

Question

In the following, the element with the highest ionization energy is (a) \([\mathrm{Ne}] 3 \mathrm{~s}^{2} 3 \mathrm{p}^{1}\) (b) \([\mathrm{Ne}] 3 \mathrm{~s}^{2} 3 \mathrm{p}^{3}\) (c) \([\mathrm{Ne}] 3 \mathrm{~s}^{2} 3 \mathrm{p}^{2}\) (d) \([\mathrm{Ne}] 3 \mathrm{~s}^{2} 3 \mathrm{p} 4\)

Step-by-Step Solution

Answer

Sulfur (d) has the highest ionization energy.

Step 1: Understand Ionization Energy

Ionization energy is the amount of energy required to remove an electron from a gaseous atom or ion. It generally increases across a period from left to right on the periodic table due to increasing nuclear charge and decreasing atomic radius.

Step 2: Review Electron Configurations

Identify the elements based on the given electron configurations: - (a) represents Aluminum (Al), - (b) represents Phosphorus (P), - (c) represents Silicon (Si), - (d) represents Sulfur (S).

Step 3: Identify Periodic Trends

Within a period, ionization energy typically increases as we move from left to right. Thus, the element further right generally has a higher ionization energy than those to the left within the same period.

Step 4: Compare and Conclude

Among the configurations provided, Sulfur (d) is the furthest to the right on the periodic table and is expected to have the highest ionization energy among the options.

Key Concepts

Periodic Trends

Periodic trends are patterns observed across periods and down groups on the periodic table. One of the most notable trends is the increase in ionization energy as you move across a period from left to right. This is due to the increasing nuclear charge; as more protons are added to the nucleus, electrons are held more tightly by the stronger attraction. Another key trend is the atomic radius, which typically decreases across a period and increases as you move down a group.
This decreased atomic radius contributes to the increase in ionization energy across a period because the outermost electrons are closer to the nucleus and more strongly attracted to it. As a result, more energy is required to remove an electron, hence the increase in ionization energy.

• Ionization energy increases across a period.
• Atomic radius decreases from left to right.
• Nuclear charge increases across a period.
• Electrons are more tightly bound as a result.

Electron Configuration

Electron configuration is crucial for understanding the behavior of elements, as it describes the distribution of electrons around an atom’s nucleus. It explains how electrons fill up various atomic orbitals in order of increasing energy levels. These configurations determine an element's chemical properties and its place in the periodic table. For instance, the given electron configurations help identify the elements Aluminum (Al), Phosphorus (P), Silicon (Si), and Sulfur (S).
Let's break down why electron configuration is essential:

• Reflects how electron shells and subshells are filled.
• Helps predict chemical bonding behavior.
• Aids in identifying elements based on unique configurations.
• Defines an element's position on the periodic table.

Understanding electron configuration also assists in explaining periodic trends, such as ionization energy, as these configurations indicate how strongly electrons are bound to an atom.

Periodic Table

The periodic table is a powerful tool that organizes elements by increasing atomic number and groups them by similar properties. This unique arrangement reflects electron configurations and recurring chemical properties. The table helps predict element behavior, such as reactivity and ionization energy trends.
The structure of the periodic table is designed to highlight periodic trends:

• Elements are organized in periods (rows) and groups (columns).
• Elements in the same group have similar electron configurations and chemical properties.
• Periods illustrate repeating patterns of properties like atomic size and ionization energy.
• The table separates metals, nonmetals, and metalloids in a visually informative way.

Understanding the periodic table's layout and how elements are positioned provides critical insight into trends like ionization energy.
This understanding is fundamental when trying to predict which elements have the highest ionization energies, as was the case with the configurations identifying Sulfur (S) as the element with the highest ionization energy in the original exercise.