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) \(14.6,14.6\)

Short answer

(a) 14.6, 13.6

Step-by-step solution

Understand the Concept

The first ionization energy is the energy required to remove the outermost electron from a neutral atom in its gaseous state. For elements in the same period, ionization energy generally increases from left to right across the periodic table.

Consider Electronic Configurations

Nitrogen has an electronic configuration of 1s² 2s² 2p³, while oxygen has 1s² 2s² 2p⁴. Nitrogen's half-filled p subshell makes it especially stable and hence requires more energy to remove an electron compared to oxygen.

Determine Ionization Energies

Based on the electronic configurations, nitrogen should have a higher first ionization energy than oxygen due to its extra stability from a half-filled p subshell. Thus, nitrogen > oxygen in ionization energy.

Choose the Correct Option

Given the options, (a) 14.6, 13.6 is consistent with our conclusion that nitrogen's first ionization energy is higher than oxygen's.

Key concepts

Electronic Configuration

To grasp the concept of ionization energy, understanding electronic configurations is key. Electronic configuration refers to the distribution of electrons in an atom's orbitals. Each element's electrons are arranged in various energy levels or shells, famously memorized as the configurations starting from 1s to higher levels like 2p. Nitrogen, for example, boasts an electronic configuration of 1s² 2s² 2p³. This notation means that nitrogen has two electrons in the 1s orbital, two in the 2s orbital, and three in the 2p orbital. On the other hand, oxygen, with its electronic configuration of 1s² 2s² 2p⁴, has an additional electron in the 2p orbital. This difference might seem minor, but it actually influences their respective ionization energies. The stability provided by nitrogen's half-filled 2p subshell contributes significantly to its higher ionization energy compared to oxygen.

Periodic Table Trends

The periodic table is not just a list of elements; it helps predict an element's chemical behavior. One of the trends observable in the periodic table is ionization energy, which generally increases across a period from left to right. This trend occurs because the effective nuclear charge increases across a period, meaning the nucleus pulls on its electrons more strongly, requiring more energy to remove an outer electron. However, certain exceptions arise due to electronic configurations, like the case with nitrogen and oxygen. Understanding these trends and the electron configurations helps students predict properties like ionization energy, especially when evaluating elements within the same period.

First Ionization Energy

The first ionization energy is a fundamental concept in chemistry. It is defined as the energy needed to remove the most loosely bound electron of an isolated gaseous atom to form a cation. This energy quantifies an atom’s stability and reluctance to lose an electron. Atoms with higher ionization energies require more energy to remove an electron, signifying that they hold onto their electrons more strongly. Ionization energy is not just a number; it reflects the interactions between the nucleus and the electrons, influenced by factors like electron shielding and subshell configurations. These attributes play a pivotal role when comparing elements like nitrogen and oxygen.

Nitrogen Ionization Energy

Nitrogen's first ionization energy is higher than that of oxygen due to its half-filled p subshell, which grants added stability. With an electronic configuration of 1s² 2s² 2p³, nitrogen's 2p orbitals are symmetrically occupied, making it energetically favorable. This configuration's stability is why nitrogen has a first ionization energy of 14.6 eV. Despite oxygen having more protons, its extra electron disrupts the symmetrical half-filled configuration, requiring less energy to remove an electron. This anomaly teaches a poignant lesson: sometimes, periodic trends like increasing ionization energy from left to right can be overridden by factors such as subshell stability.

Oxygen Ionization Energy

Oxygen exhibits a first ionization energy of 13.6 eV, slightly lower than nitrogen's. The difference arises from the electronic configuration of oxygen, 1s² 2s² 2p⁴, which results in a paired electron within its 2p subshell. This added electron-pairing creates electron-electron repulsion, which somewhat destabilizes the atom and makes removing an electron easier than in the half-filled stable configuration of nitrogen. Analyzing oxygen's ionization energy showcases how each atom’s unique arrangement affects its chemical properties, highlighting that more protons don't always equate to higher ionization energy.