Electron Configurations and Periodicity Lec 3

Exceptions to the Aufbau Principle

The Aufbau Principle states that electrons fill atomic orbitals in the order of increasing energy levels; however, there are notable exceptions among certain transition and post-transition metals due to their unique electron configurations that lead to increased stability.

Groups with Exceptions

Group 6B: This group begins with Chromium (Cr) and includes other transition metals such as Molybdenum (Mo) and Tungsten (W). These elements exhibit unique configurations to maximize the half-filled and fully filled subshell stability, which is fundamentally more favorable in terms of energy.
Group 1B: Starting with Copper (Cu), this group includes Silver (Ag) and Gold (Au). These elements similarly rearrange their electron counts to favor stability.

Electron Configuration of Copper

Using the Aufbau Principle for Copper as a case study:

Noble Gas Core: Argon (Ar) serves as the noble gas core, encapsulating the initial 18 electrons, leading to a basis of configuration.
Incorrect Initial Configuration: The traditional initial electron configuration is denoted as 4s² 3d⁹, which does not accurately reflect the observed properties of Copper.
Correction: To enhance stability, the electron configuration is corrected by promoting one electron from the 4s orbital to the 3d orbital:
- Orbital Diagram for Cu:
  - 4s orbital: 1 electron
  - 3d orbital: 10 electrons (resulting in a more stable configuration).
Correct Configuration: The correct electron configuration for Copper is 4s¹ 3d¹⁰, which provides heightened stability by optimizing electron arrangements in d orbitals.

Ground State vs. Excited State

Ground State: The arrangement of electrons in their lowest energy configuration, which ensures maximum stability. For Copper, this is achieved post-electron promotion.
Excited State: Electrons are promoted to higher energy levels without achieving a stable lowering of lower energy levels:
- Example: Phosphorus demonstrates this distinctly:
  - Ground State: 1s² 2s² 2p⁶ 3s² 3p³
  - Excited State: 1s² 2s² 2p⁶ 3s² 3p² 4s¹, in this case resulting in a configuration that is inherently less stable due to inefficient filling of lower energy levels.

Magnetic Properties

Paramagnetic: These materials contain unpaired electrons and exhibit a weak attraction to magnetic fields, valuable in applications like MRI technology.
Diamagnetic: These substances possess all electrons paired and show no attraction to magnetic fields, including noble gases and alkaline earth metals (Group 2A).

Periodic Properties: Atomic Radius

Trends in Atomic Radius: The atomic radius shows a discernible trend across the periodic table:
- Increases Down a Group: For example, Cs > Rb > K > Na > Li, as additional electron shells are added, increasing the radius.
- Decreases Across a Period: For instance, F < Ne < Na, attributed to the effective nuclear charge pulling electrons closer to the nucleus.
Effective Nuclear Charge: More protons result in stronger attractions to electrons, shrinking atomic size as one moves left to right across a period. For example, Lithium (Li) with 3 protons has less pull compared to Fluorine (F) with 9 protons due to increased shielding by inner electrons.

Ranking Atomic Radii

Example Elements: Analyzing elements such as Sulfur (S), Oxygen (O), Selenium (Se), and Arsenic (As):
- Trend: The increasing size down the group is observed as: Se > S > O.
- Cross-Period Comparison: Arsenic (As) is greater in size than Selenium (Se) due to the progressive effects of periodic trends. Hence, the final ranking based on atomic radii size is: Arsenic (largest) > Selenium > Sulfur > Oxygen (smallest).

Upcoming Topics

The next focus will pivot to the overview of ionization energy and electron affinity trends, exploring how these relate to the stability of elements.
An introduction to the periodic properties of main group elements will supplement the foundational knowledge obtained through these discussions.