Problem 92 - Introductory Chemistry Atoms First

Question

Write electron configurations for \(\mathrm{O}, \mathrm{O}^{2+}\), and \(\mathrm{O}^{2-}\). Which form would you expect to find in most compounds of oxygen? Why?

Step-by-Step Solution

Answer

The electron configurations for oxygen (O), oxygen cation (O²⁺), and oxygen anion (O²⁻) are: - O: 1s^2 2s^2 2p^4 - O²⁺: 1s^2 2s^2 2p^2 - O²⁻: 1s^2 2s^2 2p^6 The most common form of oxygen in compounds is the anion, O²⁻, because gaining 2 electrons results in a stable, full outer electron shell (2p^6), resembling the noble gas neon. This stability enables O²⁻ to readily form chemical bonds with other elements through ionic bonding.

Step 1: Identify the atomic number of oxygen

The atomic number of oxygen is 8, which means it has 8 protons and 8 electrons in its neutral form.

Step 2: Write the electron configuration for the neutral oxygen atom (O)

To write the electron configuration for O, we will follow the Aufbau principle and the Pauli exclusion principle, filling in orbitals from lower to higher energy levels: 1s^2 2s^2 2p^4 We have filled the 1s, 2s, and 2p orbitals so that the sum of electrons is 8, the total number of electrons in a neutral oxygen atom.

Step 3: Write the electron configuration for the oxygen cation with a 2+ charge (O²⁺)

For O²⁺, we need to remove two electrons from the electron configuration of the neutral oxygen atom. To do this, we will remove the electrons from the highest energy orbital, which is the 2p orbital. Oxygen cation (O²⁺) electron configuration: 1s^2 2s^2 2p^2

Step 4: Write the electron configuration for the oxygen anion with a 2- charge (O²⁻)

For O²⁻, we need to add two electrons to the electron configuration of the neutral oxygen atom. We will add the electrons to the highest energy orbital that is not yet full, which is the 2p orbital. Oxygen anion (O²⁻) electron configuration: 1s^2 2s^2 2p^6

Step 5: Identify the most common form for oxygen in compounds and explain why

The most common form of oxygen in compounds is the anion, O²⁻. This is because when oxygen gains 2 electrons, it becomes more stable, having a full outer electron shell (2p^6). In this configuration, the oxygen anion has the stable electron configuration of a noble gas – neon, which allows it to readily form chemical bonds with other elements in a process called ionic bonding.

Key Concepts

Oxygen Atom

Oxygen is a crucial element in both the environment and living organisms. Regularly found in water and organic compounds, the oxygen atom plays a vital role in life on Earth. The atomic number of oxygen is 8, which means it has 8 protons in its nucleus and, in a neutral state, 8 electrons. These electrons are configured in various shells around the nucleus, influencing the element’s chemical properties.
In its neutral form, the electron configuration of an oxygen atom is written as:

• 1s² 2s² 2p⁴

This means that the first shell holds two electrons in the 1s orbital, the second shell holds two electrons in the 2s orbital, and the remaining four electrons are in the 2p orbitals. The distribution of these electrons is vital for the atom's chemical reactivity.

Cation and Anion

Ions are formed when atoms gain or lose electrons, leading to a charge imbalance. When an atom loses electrons, it becomes a cation with a positive charge. Conversely, when it gains electrons, it forms an anion, which has a negative charge.
For oxygen:

• Cation (O²⁺): This occurs when oxygen loses two electrons, resulting in a 2+ charge. The electron configuration becomes 1s² 2s² 2p² due to the loss.
• Anion (O²⁻): This happens when oxygen gains two additional electrons, leading to a 2− charge. The electron configuration changes to 1s² 2s² 2p⁶, similar to the noble gas, neon.

Anions like O²⁻ are common and stable because achieving a full outer shell mimics the configuration of noble gases, which are known for their stability.

Electron Orbitals

Electron orbitals describe regions in an atom where electrons are likely to be found. These orbitals are filled with electrons in a specific order, following rules known as the Aufbau principle, Hund's rule, and the Pauli exclusion principle.
Oxygen's electrons are located in the:

• 1s orbital: the closest to the nucleus and holds 2 electrons.
• 2s orbital: next in line and can also accommodate 2 electrons.
• 2p orbitals: the highest energy level for electrons in oxygen, allowing up to 6 electrons. In its basic state, they hold 4 electrons.

Electron configuration dictates how an atom will interact chemically. When we alter the configuration by adding or removing electrons, it forms ions, affecting the atom's reactivity and bonding characteristics.

Atomic Number

The atomic number is fundamental to defining an element. It represents the number of protons in the nucleus and is unique for each element. Oxygen has an atomic number of 8, identifying it unequivocally among other elements.
Understanding an element’s atomic number helps determine its position in the periodic table and predicts its chemical behavior.

• The atomic number equals the number of electrons in a neutral atom.
• It affects the element's electron configuration and reactivity.

In ions, the electron number can change, but the atomic number—representing protons—remains constant, ensuring the element keeps its identity.

Ionic Bonding

Ionic bonding is a chemical bond formed between two ions with opposite charges. This type of bonding occurs when one atom donates electrons to another, resulting in a cation and an anion.
For oxygen, which commonly forms an O²⁻ anion, ionic bonding is typical when it reacts with elements that lose electrons (such as metals).

• The positive cation and the negative anion attract each other, creating a stable compound.
• This process often results in the formation of salts, like sodium chloride (table salt).

Ionic bonds give rise to compounds with high melting and boiling points, and they often dissolve well in water, conducting electricity in solution. Understanding ionic bonding provides insight into the structure and behavior of various substances.