Problem 46 - 43 Years JEE ADVANCED (1978-2020) + JEE MAIN Chapterwise & Topicwise Solved Papers Chemistry

Question

The element with the highest first ionization potential is (a) boron (b) carbon (c) nitrogen (d) oxygen

Step-by-Step Solution

Answer

Nitrogen (c) has the highest first ionization potential.

Step 1: Understand Ionization Potential

The first ionization potential is the energy required to remove the outermost electron from a gaseous atom. Elements with a more stable electronic configuration generally have a higher ionization potential.

Step 2: Recognize Periodic Trends

As we move across a period in the periodic table from left to right, the ionization potential generally increases. This is because the atomic radii decrease and the effective nuclear charge increases, holding electrons more tightly.

Step 3: Analyze Elements in the Group

Within the provided elements boron (B), carbon (C), nitrogen (N), and oxygen (O), observe their positions in the periodic table. They belong to the same period, specifically period 2.

Step 4: Compare Electronic Configurations

Review the electronic configurations of the elements: \( \text{B: } [1s^2 2s^2 2p^1], \text{C: } [1s^2 2s^2 2p^2], \text{N: } [1s^2 2s^2 2p^3], \text{O: } [1s^2 2s^2 2p^4] \). Nitrogen has a half-filled \(2p\) subshell, which is a particularly stable configuration.

Step 5: Determine the Highest Ionization Potential

Considering the periodic trend and the stability of the half-filled \(2p\) subshell, nitrogen, with an atomic configuration of \(1s^2 2s^2 2p^3\), has the highest ionization potential among the options given.

Key Concepts

Periodic Table Trends

The periodic table is an incredible tool for understanding the properties of elements, including ionization potential. As you move across a period from left to right, several things happen.

• The atomic number increases, adding more protons to the nucleus, which in turn increases the nuclear charge. This means electrons are pulled closer to the nucleus.
• The atomic size, or atomic radii, decreases because of this increased attraction between the nucleus and the electrons.
• The ionization potential generally increases. This trend occurs because electrons are held more tightly and more energy is needed to remove the outermost electron.

For example, in a single period, such as period 2, nitrogen has a higher ionization potential than carbon or boron. This is due to the tighter hold the nucleus has on its electrons as we move across the period.

Electronic Configuration

Electronic configurations provide a roadmap of how electrons are arranged in an atom. Each electron has a specific energy level, described in terms of shells and subshells. Here’s a brief overview:

• Electrons fill orbitals in a way that minimizes the energy of the atom. This is why we follow the "Aufbau principle," filling from lower to higher energy levels.
• For instance, **boron** has an electronic configuration of \(1s^2 2s^2 2p^1\), while **carbon** is \(1s^2 2s^2 2p^2\).
• **Nitrogen** is \(1s^2 2s^2 2p^3\), and **oxygen** is \(1s^2 2s^2 2p^4\).

These configurations tell us a lot about how atoms interact. An electron's position within an atom affects its chemical bonding and properties like ionization potential, as explained through periodic trends.Understanding these configurations helps us see why certain elements, like nitrogen, have particularly high or low ionization potentials due to more stable electronic arrangements.

Stable Subshell Configurations

The stability of an atom's electronic structure is greatly influenced by its subshell configurations. Stable configurations usually result in higher ionization potentials. Here are a few key points:

• A half-filled or completely filled subshell is often more stable. In our example, **nitrogen** (\(1s^2 2s^2 2p^3\)) has a half-filled \(2p\) subshell.
• This half-filled status gives nitrogen extra stability compared to elements like oxygen, which has an additional electron making it \(1s^2 2s^2 2p^4\). Adding more electrons after the half-filled stage often causes repulsion, slightly lowering stability.
• Stable configurations resist ionization, meaning they require more energy to remove an electron.

Thus, even though oxygen is to the right of nitrogen in the periodic table, nitrogen's half-filled \(2p\) subshell makes its configuration remarkably robust, leading to a higher ionization potential. This stability is key in understanding why it requires more energy to ionize nitrogen.