Question

Use orbital diagrams to illustrate what happens when an oxygen atom gains two electrons. On the basis of what you have learned about electron affinities and electron configurations, why is it extremely difficult to place a third electron on the oxygen atom?

Short answer

When an oxygen atom gains two electrons, they fill the remaining two spots in the 2p orbitals, creating a full valence shell. Adding a third electron is difficult due to the high electron-electron repulsion and the need to overcome the oxygen atom's high effective nuclear charge.

Step-by-step solution

Identify Oxygen's Electron Configuration

First, determine the ground-state electron configuration of an oxygen atom. Oxygen has an atomic number of 8, which means it has 8 electrons. The electron configuration for oxygen is 1s² 2s² 2p⁴. This shows that the first two energy levels, the 1s and 2s orbitals, are fully filled, while the 2p orbital has four electrons out of a maximum of six.

Draw the Orbital Diagram for Oxygen

Now, draw the orbital diagram for the ground-state electron configuration of oxygen. Each orbital is represented by a box, and each electron by an arrow. The arrows are drawn in opposite directions to represent electron spin. The filled orbitals (1s and 2s) will have two arrows pointing in opposite directions in each box, and for the 2p orbital, you should have three boxes with one arrow in each for the first three electrons, and a fourth box with an opposite arrow to represent the fourth electron in the 2p orbital.

Add Two Electrons to the Orbital Diagram

When an oxygen atom gains two electrons, they will go into the next available spots to minimize electron repulsion, following Hund's Rule. Thus, add one electron (arrow) each to the two empty 2p orbital boxes. The final orbital diagram will have the 2p level fully occupied, with all three boxes containing two electrons each.

Describe the Difficulty of Adding a Third Electron

When attempting to add a third electron, the electron would have to enter the next higher energy level, which would be the 3s orbital. This is energetically unfavorable because oxygen has high effective nuclear charge and electron-electron repulsion within a relatively small atom. Consequently, adding another electron would experience significant repulsion and would require a great amount of energy to overcome the electrostatic forces. This is why the electron affinity for the third electron is positive, indicating an endothermic process.

Key concepts

Orbital Diagrams

Orbital diagrams are a visual way to represent the arrangement of electrons in an atom's orbitals. They use boxes (or lines) to represent orbitals and arrows to depict electrons, with the direction of the arrow indicating the electron's spin. One of the essential rules followed while filling these diagrams is the Pauli Exclusion Principle, which states that no two electrons in an atom can have the same set of four quantum numbers; hence, an orbital can hold only two electrons with opposite spins.

For oxygen, which has eight electrons, the first two fill the 1s orbital, the next two are in the 2s orbital, and the last four are distributed among the three 2p orbitals. The correct filling of these orbitals is crucial to understanding how the atom behaves when it gains or loses electrons, such as its capacity to form bonds with other atoms.

Electron Affinity

Electron affinity refers to the amount of energy released when an electron is added to a neutral atom in the gaseous state to form a negative ion. It provides insight into the atom's tendency to accept additional electrons. In general, atoms release energy when they acquire an electron, which is characteristic of a negative electron affinity value. However, as in the case of adding a third electron to oxygen, the electron affinity becomes positive, indicating that energy must be supplied for the addition to occur.

The difficulty in placing a third electron on the oxygen atom arises from the atomic structure that is already stable with eight electrons. Adding another electron means overcoming significant electron-electron repulsion and the electrostatic attraction between electrons and the nucleus, leading to a less stable and energetically unfavorable configuration.

Hund's Rule

Hund's Rule plays a pivotal role in predicting the ground-state electron configurations of atoms. It states that electrons will occupy separate orbitals of the same energy level before they pair up. This behavior minimizes electron-electron repulsion since electrons are negatively charged and repel each other. They fill orbitals in a way that maximizes the number of half-filled orbitals.

When applying Hund's Rule to oxygen's 2p orbitals, one electron enters each orbital until all are half-filled before any pairing occurs. This leads to a configuration with two unpaired electrons in the 2p orbital, which can then pair up as two additional electrons are added. Hund's Rule explains why the oxygen atom tends to gain exactly two electrons to achieve a full 2p sublevel, resulting in a stable configuration.

Ground-State Electron Configuration

The ground-state electron configuration of an atom is the arrangement of electrons in the lowest energy levels possible. For oxygen with eight electrons, the configuration is represented as 1s² 2s² 2p⁴. This notation describes the filled 1s and 2s orbitals and the partially filled 2p orbital, which has four electrons out of a maximum of six.

The ground-state configuration is vital for understanding an atom's reactivity, the types of bonds it can form, and how it interacts with other atoms. Oxygen, due to its electron configuration, is highly reactive and readily forms covalent bonds by sharing its 2p electrons to achieve a stable noble gas configuration, often by attracting electrons from other atoms to fill its 2p orbitals.