Electron Configurations and Periodic Table Trends: Study Notes

Electron Configurations

Introduction to Electron Configurations

Electron configuration describes the arrangement of electrons in an atom's orbitals. Understanding electron configurations is essential for predicting chemical properties and reactivity.

• Abbreviations: Electron configurations can be abbreviated using noble gas notation to simplify representation.

• Mapping to Periodic Table: The periodic table structure reflects the filling order of electron orbitals.

• Formation of Ions: Electron configurations change when atoms gain or lose electrons to form ions.

• Isoelectronic Species: Atoms or ions with the same electron configuration are called isoelectronic.

Quantum Numbers and Energy Levels

Principal Energy Levels (Shells)

Electrons occupy energy levels, also known as shells, denoted by the principal quantum number n (n = 1, 2, 3, ...).

• Sublevels: Each energy level contains sublevels: s, p, d, and f, with increasing energy.

• Orbitals: Sublevels contain orbitals, which are regions where electrons are likely to be found.

• Number of Orbitals per Sublevel:

  • s: 1 orbital

  • p: 3 orbitals

  • d: 5 orbitals

  • f: 7 orbitals

• Maximum Electrons per Orbital: Each orbital can hold a maximum of 2 electrons.

Example Calculation: For n = 2 (second energy level): - 2s sublevel: 1 orbital × 2 electrons = 2 electrons - 2p sublevel: 3 orbitals × 2 electrons = 6 electrons - Total electrons in n = 2: 8 electrons

$\text{Maximum electrons in a shell} = \text{Number of orbitals} \times 2$

Rules for Electron Configuration

Aufbau Principle

Electrons fill the lowest available energy orbitals first before occupying higher energy levels.

• Order of Filling: The order is determined by increasing energy, generally following the sequence: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc.

Pauli Exclusion Principle

No two electrons in the same atom can have the same set of quantum numbers. In practice, this means:

• Each orbital can hold a maximum of 2 electrons, and they must have opposite spins.

Hund's Rule

Electrons will fill degenerate (equal energy) orbitals singly before pairing up. This minimizes electron repulsion and increases stability.

• For example, in the p sublevel (three orbitals), one electron enters each orbital before any orbital gets a second electron.

Orbital Diagrams and Examples

Orbital Diagrams for Selected Atoms

Orbital diagrams visually represent electron configurations, showing electrons as arrows in boxes (orbitals).

Atom	Electron Configuration	Orbital Diagram	Total Electrons
Hydrogen (H)	1s1	1 box with 1 arrow	1
Helium (He)	1s2	1 box with 2 arrows (opposite spins)	2
Lithium (Li)	1s2 2s1	1s: 2 arrows; 2s: 1 arrow	3
Magnesium (Mg)	1s2 2s2 2p6 3s2	1s: 2 arrows; 2s: 2 arrows; 2p: 6 arrows; 3s: 2 arrows	12
Boron (B)	1s2 2s2 2p1	1s: 2 arrows; 2s: 2 arrows; 2p: 1 arrow	5
Carbon (C)	1s2 2s2 2p2	1s: 2 arrows; 2s: 2 arrows; 2p: 2 arrows (in separate orbitals)	6
Arsenic (As)	1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3	All lower orbitals filled; 4p: 3 arrows (Hund's Rule)	33

Periodic Table Trends

Introduction to Periodic Trends

The arrangement of electrons in atoms explains many periodic trends, such as atomic radius, ionization energy, and electron affinity.

• Atomic Radius: Generally decreases across a period and increases down a group.

• Ionization Energy: Increases across a period and decreases down a group.

• Electron Affinity: Tends to become more negative across a period.

Example: Elements in the same group have similar valence electron configurations, leading to similar chemical properties.

Formation of Ions and Isoelectronic Species

Formation of Ions

Atoms gain or lose electrons to achieve a stable electron configuration, often resembling the nearest noble gas.

• Cations: Formed by loss of electrons (e.g., Na+).

• Anions: Formed by gain of electrons (e.g., Cl-).

Isoelectronic Species

Isoelectronic species have identical electron configurations but may be different elements or ions.

• Example: Ne (Neon), Na+, and F- all have the electron configuration 1s2 2s2 2p6.

Additional info: These notes expand on the original class notes by providing definitions, examples, and context for each topic, ensuring a self-contained study guide suitable for introductory chemistry students.