Arrangement of Electrons in the Energy Levels of Atom

The arrangement of electrons in the energy levels of an atom is described by the electron configuration. This configuration is determined by a set of principles that guide how electrons fill available atomic orbitals. Here’s a detailed look at these principles and how they determine electron arrangement:

Key Principles

1. Aufbau Principle:

   • Electrons fill orbitals starting with the lowest energy level before moving to higher levels.
   • The order of filling is determined by the increasing energy of the orbitals, which can be remembered using the Aufbau diagram or the sequence derived from the periodic table.

2. Pauli Exclusion Principle:

   • No two electrons in an atom can have the same set of four quantum numbers.
   • This means each orbital can hold a maximum of two electrons with opposite spins.

3. Hund's Rule:

   • When electrons occupy orbitals of equal energy (degenerate orbitals), one electron enters each orbital until all orbitals contain one electron with parallel spins.
   • Only after all degenerate orbitals are half-filled do electrons pair up.

Energy Levels and Sublevels

Electrons are arranged in energy levels (shells) denoted by the principal quantum number $𝑛$. Each energy level contains sublevels (subshells) designated by the azimuthal quantum number $𝑙$:

• n = 1: 1s
• n = 2: 2s, 2p
• n = 3: 3s, 3p, 3d
• n = 4: 4s, 4p, 4d, 4f

Each subshell contains a specific number of orbitals, and each orbital can hold two electrons:

• s sublevel: 1 orbital (2 electrons)
• p sublevel: 3 orbitals (6 electrons)
• d sublevel: 5 orbitals (10 electrons)
• f sublevel: 7 orbitals (14 electrons)

Electron Configuration Notation

Electron configurations are written using the notation $𝑛{\mathrm{\ell }}^{𝑥}$, where $𝑛$ is the principal quantum number, $\mathrm{\ell }$ is the type of orbital (s, p, d, f), and $𝑥$ is the number of electrons in that subshell.

Examples

1. Hydrogen (H):

   • Atomic number 1
   • Configuration: $1{𝑠}^1$

2. Helium (He):

   • Atomic number 2
   • Configuration: $1{𝑠}^2$

3. Carbon (C):

   • Atomic number 6
   • Configuration: $1{𝑠}^{2}2{𝑠}^{2}2{𝑝}^2$

4. Neon (Ne):

   • Atomic number 10
   • Configuration: $1{𝑠}^{2}2{𝑠}^{2}2{𝑝}^6$

5. Sodium (Na):

   • Atomic number 11
   • Configuration: $1{𝑠}^{2}2{𝑠}^{2}2{𝑝}^{6}3{𝑠}^1$

Electron Configuration for Transition Metals and Beyond

For elements beyond the second period, the order of filling can be more complex due to the relative energies of the 3d and 4s orbitals. For example:

• Iron (Fe):
  • Atomic number 26
  • Configuration: $[𝐴𝑟]3{𝑑}^{6}4{𝑠}^2$

Periodic Table and Electron Configuration

The periodic table is structured in such a way that elements in the same group have similar electron configurations in their outermost shells, leading to similar chemical properties. The periodic trends and the filling order can be visualized using the periodic table's blocks:

• s-block: Groups 1 and 2 (and Helium)
• p-block: Groups 13 to 18
• d-block: Transition metals (Groups 3 to 12)
• f-block: Lanthanides and Actinides

Conclusion

Understanding the arrangement of electrons within an atom's energy levels is fundamental to the study of chemistry and physics. It explains the periodic properties of elements, their reactivity, and the formation of chemical bonds. For more detailed and specific configurations, consulting the periodic table and the rules mentioned above is essential.