Problem 54 - A Complete Resource Book in Chemistry for JEE Main

Question

The first ionization energy in electron volts of nitrogen and oxygen atoms are respectively given by (a) \(14.6,13.6\) (b) \(13.6,14.6\) (c) \(13.6,13.6\) (d) \(14.6,14.6\)

Step-by-Step Solution

Answer

(a) 14.6, 13.6

Step 1: Understanding Ionization Energy

Ionization energy is the minimum amount of energy required to remove the outermost electron from a neutral atom in its gaseous state. The values provided are in electron volts (eV) and refer to the energy required to ionize nitrogen and oxygen atoms.

Step 2: Review Periodic Trends

Ionization energy tends to increase across a period from left to right on the periodic table. Nitrogen is to the left of oxygen in the periodic table, which generally indicates that nitrogen should have a lower ionization energy than oxygen. However, nitrogen has a half-filled p subshell, making it more stable and thus requiring more energy to remove an electron compared to oxygen.

Step 3: Compare Nitrogen and Oxygen

In this specific case, despite the usual trend, nitrogen has a higher ionization energy than oxygen due to its stable half-filled electronic configuration. Thus, nitrogen's first ionization energy is higher at 14.6 eV compared to oxygen's 13.6 eV.

Step 4: Select Correct Option

Given the ionization energies, select the option where nitrogen's ionization energy is 14.6 eV and oxygen's is 13.6 eV. Therefore, the correct choice is (a) 14.6, 13.6.

Key Concepts

Electron Configuration

Electron configuration describes how electrons are distributed in an atom’s atomic orbitals. Each element has a unique electron configuration, which determines its chemical properties and stability.

• Electrons fill orbitals starting from the lowest energy level to higher levels, a process guided by the Aufbau principle.
• The configuration is denoted using the notation such as 1s, 2s, 2p, indicating different energy levels and sub-levels where electrons are found.

Understanding electron configuration is crucial because it helps explain the chemical behavior of elements. The way electrons are arranged dictates how atoms interact with each other. For instance, elements with half-filled or fully filled subshells, like nitrogen with its half-filled 2p subshell, often display unique stability that influences their ionization energy.

Periodic Trends

Periodic trends refer to patterns observed in the properties of elements across different periods and groups on the periodic table. Knowing these trends helps predict the properties of unknown elements.

• Ionization energy, electronegativity, and atomic radius are some of the key properties demonstrating periodic trends.
• Generally, ionization energy increases across a period (left to right) and decreases down a group (top to bottom).

These trends occur due to variations in electron configurations and the effective nuclear charge (the net positive charge experienced by an electron). For example, moving across a period, electrons are added to the same energy level while the nucleus' positive charge increases, leading to a higher attraction for electrons and a rise in ionization energy.

Nitrogen Ionization Energy

Nitrogen's ionization energy is an interesting anomaly in periodic trends. Despite being to the left of oxygen on the periodic table, nitrogen has a higher first ionization energy.

• Nitrogen has the electron configuration of 1s² 2s² 2p³, giving it a stable half-filled p subshell.
• This half-filled stability means more energy is required to remove an electron from nitrogen compared to the trend expected purely based on its position relative to oxygen.

This unique stability associated with its half-filled subshell is a key reason why nitrogen's ionization energy is measured at 14.6 eV. This value is higher than that of oxygen, opposing the general trend but highlighting the importance of electron configuration in determining ionization energy.

Oxygen Ionization Energy

Oxygen, with an electron configuration of 1s² 2s² 2p⁴, follows an interesting pattern of ionization energy in comparison to nitrogen due to its electron arrangement.

• Oxygen needs less energy to remove an electron because it can relieve electron-electron repulsion by transitioning from a 2p⁴ to a more stable configuration after ionization.
• This slightly decreased stability compared to nitrogen makes oxygen's first ionization energy lower, valued at 13.6 eV.

Though typically elements to the right in a period exhibit higher ionization energies, oxygen’s particular electron arrangement causes a lower energy requirement compared to nitrogen. This deviation underscores the importance of understanding individual electron configurations alongside periodic trends when analyzing chemical properties.