Electron Configuration (AQA A-Level Chemistry): Revision Notes

1.1.5 Electron Configuration

Key Concepts

• Electron shells: Energy levels where electrons are arranged around the nucleus.
• Sub-shells: Orbitals within each shell that can hold a certain number of electrons. These are labelled s, p, d, and f.

Electron Configurations up to $Z$ = 36

For the first 36 elements (atomic number $Z$ = 36), electrons occupy the s, p, and d sub-shells in a specific order. This section will explain how to write electron configurations for these elements, according to the AQA A-level Chemistry specification.

Aufbau Principle

The Aufbau Principle dictates that electrons fill orbitals starting with the lowest energy levels first, moving upwards. This order can be memorised using the following sequence of orbitals:

$$1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p$$

Note that the 4s orbital fills before 3d because it is lower in energy. Therefore, configurations such as [Ar] 4s² 3d⁶ are common for transition metals like iron (Fe).

Electron Configuration of Elements

Here is a breakdown of how electrons fill the orbitals for the elements up to Z = 36 (Krypton, Kr):

Writing Electron Configurations for Atoms and Ions

Notation:

Write the number of electrons in each sub-shell as a superscript. For example, oxygen (O), with 8 electrons, is written as:

1s² 2s² 2p⁴.

Ions:

When writing the electron configuration for ions, remember that electrons are lost or gained from the outermost sub-shell. For example, Mg²⁺ (magnesium ion) loses two electrons from the 3s orbital:

1s² 2s² 2p⁶.

Noble Gas Shorthand:

To simplify long electron configurations, noble gas symbols in square brackets can be used. For instance, the configuration for potassium (K) can be written as:

[Ar] 4s¹

(where [Ar] = 1s² 2s² 2p⁶ 3s² 3p⁶).

Sub-shell Notations

s, p, d-block Elements

• s-block elements: Elements where the outermost electrons are in an s-orbital. This includes Group 1 and Group 2 elements like hydrogen, lithium, and magnesium.
• p-block elements: These elements have their outermost electrons in a p-orbital, found in Groups 13 to 18 (e.g., carbon, nitrogen, oxygen).
• d-block elements: Elements where electrons are filling the d-orbitals. These are the transition metals found in the centre of the periodic table (e.g., iron, copper).

Important Points

1. The 4s orbital fills before the 3d orbital, but when removing electrons (for ions), electrons are lost from the 4s orbital first.
2. Hund's Rule: Electrons occupy orbitals singly as much as possible before pairing up.
3. Pauli Exclusion Principle: No two electrons in the same orbital can have the same spin.

Examples of Electron Configurations