CHEM 1411 Concept Reviews: Periodic Properties of the Elements Part I

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Electron configuration: a way of representing the locations and energies of all the electrons within an atom. Below are several rules and principles that must be considered when determining electron configuration.

Aufbau Principle: Electrons will always occupy lower energy orbitals before higher energy orbitals.

Hund’s Rule: Electrons will fill orbitals of equal energy (degenerate orbitals) with one electron with the same spin (parallel spin) in each orbital before any of those orbitals will gain a second electron.

Pauli Exclusion Principle: No two electrons can have the same set of 4 quantum numbers. (Basically, this means no two electrons can be in the same orbital with the same spin.)

Writing Electron Configurations

The periodic table can be used as a roadmap for electron configuration. Each “box” on the periodic table represents a place for one electron. The different subshells are represented by different “blocks” in the periodic table; the s-block, the p-block, the d-block, and the f-block. These blocks are color coded in the diagram below.

For a complete electron configuration:

Start at the beginning (1s) and fill the orbitals in order, going left to right until you reach the end of a period and then going to the next row. Indicate each the energy level, the subshell being filled, and how many electrons are filling it as you pass through each subshell. Represent the symbols in the following manner:

CHEM 1411 Ch 8-1 img 1.png

Continue filling subshells until you reach the element for which you are determining the electron configuration. Remember that the f-orbitals are actually in between the 6s & 5d and the 7s & 6d orbitals. The subshell filling guide below on the right can help you check to make sure that you have filled the orbitals in the correct order. Also note that subshells drop 1 energy level when you pass into the d-block and 2 energy levels when you pass into the f-block.

CHEM 1411 Ch 8-1 img 2.png CHEM 1411 Ch 8-1 img 3.png

Example Configurations: O: 1s² 2s² 2p⁴

Fe: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶

Exceptions to the Ordinary Filling Pattern of Electron Configuration

There are several unique configuration patterns that come up with certain elements on the periodic table. Some of the patterns are consistent and predictable. Others are seemingly more random. The most important of these are as follows:

Half-filled and completely filled subshells are much more stable than subshells of the same type which are partially filled in different ways. This stability is more significant in larger subshells such as d or f. Due to this fact, elements in the d-block that would normally be expected to have 4 or 9 electrons in the d-subshell will steal an electron from the previously filled s-subshell, giving it either 5 or 10 d-electrons, which makes it half-filled or completely filled. See examples below:

Expected Configuration Based on Filling Pattern Actual Configuration

Cr: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁴ Cr: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵

Cu: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁹ Cu: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰

After filling the 6s or 7s orbitals, a single electron will sometimes fill the d orbital that follows before dropping into the f-block. The pattern to this is seemingly random except where the half-filled & completely filled principle applies. Instead of pulling an electron from the closest s-subshell, a single electron will fill or not fill the nearest d-subshell in order to ensure that a half-filled or completely filled f-subshell is obtained.

Orbital Diagrams

Orbital Diagram: a pictorial representation of electron configuration that uses boxes to represent the orbitals of a subshell and arrows to represent electrons occupying a subshell.

The different subshells in an atom have different numbers of orbitals and different shapes according to the type of subshell. The following table summarizes this information: