Question

Write the full orbital diagram for each element. a. \(\mathrm{N}\) b. F c. Mg d. Al

Short answer

The orbital diagrams for each element are as follows: N (1s² 2s² 2p³), F (1s² 2s² 2p⁵), Mg (1s² 2s² 2p⁶ 3s²), and Al (1s² 2s² 2p⁶ 3s² 3p¹).

Step-by-step solution

Understanding Orbital Diagrams

Orbital diagrams are visual representations of the electron configurations of atoms. They show the placement of electrons in the orbitals around the nucleus of an atom. Each orbital is represented by a box or line, and each electron is represented by an arrow (↑) for one spin direction and (↓) for the opposite. The Pauli exclusion principle states that each orbital can hold a maximum of two electrons with opposite spins.

Writing the Orbital Diagram for Nitrogen (N)

Nitrogen has an atomic number of 7, implying seven electrons. Following the Aufbau principle, Hund's rule, and the Pauli exclusion principle, fill the orbitals in the order of increasing energy (1s, 2s, 2p, 3s...). For nitrogen, the orbital filling is: 1s² 2s² 2p³. Start filling from the lowest energy level (1s) moving to higher ones, and place three unpaired electrons in the three 2p orbitals to comply with Hund's rule.

Writing the Orbital Diagram for Fluorine (F)

Fluorine has an atomic number of 9, which means it has nine electrons. Following the same rules as for nitrogen, fill the orbitals in increasing energy order. The orbital filling is: 1s² 2s² 2p⁵. After filling the 1s and 2s orbitals, place a single electron in each of the three 2p orbitals, and then start pairing them up, ending with one of the 2p orbitals having a pair of electrons.

Writing the Orbital Diagram for Magnesium (Mg)

Magnesium has an atomic number of 12, indicating twelve electrons. Fill up the orbitals following the same rules: The orbital filling is 1s² 2s² 2p⁶ 3s². After the 1s, 2s, and 2p orbitals are completely filled, add the remaining two electrons to the 3s orbital.

Writing the Orbital Diagram for Aluminum (Al)

Aluminum's atomic number is 13, denoting thirteen electrons. Fill the orbitals in the same manner as the previous elements. Its orbital filling is: 1s² 2s² 2p⁶ 3s² 3p¹. After completely filling the 1s, 2s, and 2p orbitals, and placing two electrons in the 3s orbital, place the final electron in one of the 3p orbitals.

Key concepts

Electron Configurations

Understanding electron configurations is essential for grasping the behavior of atoms in different elements. Electron configurations refer to the distribution of electrons in an atom's orbitals, which are various regions around the nucleus where electrons are most likely to be found. It’s like having an address book for each electron in an atom.

Each orbital can be identified by a unique combination of quantum numbers, and electrons fill these orbitals in a manner that minimizes the atom's total energy. This means that electrons will occupy the lowest available energy level before moving to higher levels.

For example, the electron configuration for Nitrogen (N), which has seven electrons, would be written as 1s² 2s² 2p³. This indicates that two electrons fill the '1s' orbital, two fill the '2s' orbital, and the remaining three are distributed across the '2p' orbitals, reflecting the atom's most stable arrangement.

Aufbau Principle

The Aufbau principle is a fundamental guideline used to determine the electron configuration of an atom. The word 'Aufbau' is German for 'building up', and this principle essentially describes how electron orbitals are filled from the lowest to the highest energy levels.

It directs us to start filling up the electronic orbitals with the lowest energy first, and only move to higher energy levels once the lower ones are full. The order in which these orbitals are filled is represented by a specific sequence: 1s, 2s, 2p, 3s, 3p, and so on.

Aufbau Filling Order

• 1s
• 2s
• 2p
• 3s
• 3p
• 4s
• 3d
• 4p

This order reflects the relative energies of the orbitals and is an invaluable tool for writing out the electron configurations. Notably, there are exceptions to the Aufbau principle, primarily involving elements with d and f orbitals, but this pattern holds true for most simpler cases.

Hund's Rule

Hund's rule addresses the way electrons are distributed among orbitals of the same energy—known as 'degenerate' orbitals. It's kind of like a bus seating rule: if empty seats are available, it's more comfortable for passengers (electrons) to sit by themselves before sharing a seat.

According to Hund's rule, every orbital in a subshell is singly occupied with one electron before any orbital is doubly occupied. Furthermore, all electrons in singly occupied orbitals have the same spin direction. This minimizes electron repulsion, as electrons, being negatively charged, naturally repel each other, and thus makes the atom more stable.

Returning to the example of Nitrogen (N) with its electron configuration of 1s² 2s² 2p³, Hund's rule explains why the three electrons in the 2p subshell each occupy separate orbitals rather than pair up. This results in three unpaired electrons, each in a different '2p' orbital, with their spins aligned.