Electron Configurations and Periodicity Lec 1

Introduction to Electron Configurations

Understanding electron configurations is fundamental for grasping the behavior of atoms and how they interact with one another in chemical reactions. This topic connects closely with quantum numbers, which provide a framework for describing the locations and energy levels of electrons.

Key Principles and Terms

Pauli Exclusion Principle

• States that no two electrons can occupy the same quantum state simultaneously within an atom. In practical terms, this means that each of the four quantum numbers must be unique for each electron in an atom.

• For example, two electrons can share three quantum numbers (same energy state), but the fourth quantum number, which corresponds to their spin, must differ (one may be spin-up, +1/2, and the other spin-down, -1/2).

Aufbau Principle (Building Up Principle)

• Provides a method for establishing the ground state electron configuration of an atom. According to this principle, electrons populate the available atomic orbitals starting from the lowest energy levels and moving to higher levels in a systematic manner.

• The lowest energy configuration is referred to as the ground state, while an excited state occurs when electrons absorb energy and are promoted to orbitals of higher energy.

Hund's Rule

• States that when electrons are distributed among degenerate orbitals (orbitals of the same energy), one electron is placed in each orbital before any pairing occurs. This rule is crucial when drawing orbital diagrams as it minimizes electron-electron repulsion, stabilizing the atom.

Electron Shells, Subshells, and Orbitals

• Shells: Regions around the nucleus of an atom that quantify energy levels. The innermost shell (n=1) is closest to the nucleus and holds the least energy, while subsequent shells (n=2, n=3, etc.) are further out and possess higher energy levels.

• Subshells: They are further divisions of shells, designated by the letters s, p, d, and f.

  • s subshell: Can hold a maximum of 2 electrons and consists of 1 orbital.

  • p subshell: Can hold a maximum of 6 electrons and contains 3 orbitals, each oriented at right angles to each other.

  • d subshell: Holds up to 10 electrons with 5 orbitals. It has more complex shapes compared to s and p orbitals.

  • f subshell: Holds up to 14 electrons and contains 7 orbitals. Typically, f subshells are relevant for lanthanides and actinides, which will be discussed later.

Filling Orbitals

• Filling Order: Electrons fill atomic orbitals according to their increasing energy levels. This is visually demonstrated by the Aufbau principle, where energy levels must be respected when placing electrons in orbitals.

• Hund's Rule Application: When filling multiple orbitals of the same energy, each must be singly occupied with arrows pointing in the same direction representing spin orientation, before any pairing occurs.

Writing Electron Configurations

• The general format for writing electron configurations is: Shell number (n), subshell letter (s, p, d, f), followed by the number of electrons in that subshell as a superscript.

• Examples of Configurations:

  • A simple configuration may look like: 1s², 2s², 2p⁶.

  • A key point in this pattern is that after filling the 3p subshell (3p⁶), the next electrons will fill the 4s subshell rather than the 3d, because 4s is lower in energy than 3d despite being in a higher principal energy level.

Understanding Zones in the Periodic Table

• s Zone: Consists of Groups 1A (alkali metals) and 2A (alkaline earth metals), including helium, which is unique in that it behaves like a noble gas despite being placed in the s block.

• p Zone: Encompasses Groups 3A to 8A (or groups 13 to 18), containing elements that typically display a range of oxidation states.

• d Zone: Comprising the transition metals (3B to 2B or groups 3 to 12), these elements exhibit variable valencies and form colored compounds.

• f Zone: Includes the lanthanides and actinides, which are located in two rows below the main body of the periodic table. These elements are significant for their magnetic and radioactive properties.

Example: Electron Configurations Using the Periodic Table

• Zinc (Zn): Configuration: 1s², 2s², 2p⁶, 3s², 3p⁶, 4s², 3d¹⁰.

• Bromine (Br): Configuration: 1s², 2s², 2p⁶, 3s², 3p⁶, 4s², 3d¹⁰, 4p⁵.

• Antimony (Sb): Configuration: 1s², 2s², 2p⁶, 3s², 3p⁶, 4s², 3d¹⁰, 4p⁶, 5s², 4d¹⁰, 5p³.

Summary and Upcoming Topics

In the next section, discussions will delve into more complex aspects of electron configurations, particularly focusing on the f zone electron configurations and the noble gas shorthand method, which simplifies the notation for electron configurations by using the closest preceding noble gas as a reference.