3.3 Pauli exclusion principle and Hund's rules

The Pauli exclusion principle and Hund's rules are key concepts in understanding atomic structure. They explain how electrons arrange themselves in orbitals, shaping an atom's properties and behavior. These principles are crucial for grasping the periodic table and predicting chemical reactions.

These rules also impact atomic stability and reactivity. By determining electron configurations, they help us understand why some atoms are more stable or reactive than others. This knowledge is essential for predicting how elements will interact and form compounds.

Pauli exclusion principle

Fundamental principle and implications

• The Pauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers (n, l, ml, ms)
• This principle implies that each orbital can hold a maximum of two electrons with opposite spins (ms = +1/2 and -1/2)
• The Pauli exclusion principle determines the arrangement of electrons in an atom's orbitals, leading to the concept of electron configuration

Electron configuration and periodic table

• Electron configuration is the distribution of electrons among the available orbitals in an atom, following the Aufbau principle, Hund's rules, and the Pauli exclusion principle
• The Pauli exclusion principle is responsible for the structure of the periodic table and the chemical properties of elements
• Example: The electron configuration of carbon (atomic number 6) is 1s² 2s² 2p², which follows the Pauli exclusion principle and explains its tetravalent nature

Hund's rules for electron configuration

Hund's three rules

• Hund's first rule states that electrons occupy orbitals of the same energy (degenerate orbitals) singly before pairing up, with their spins aligned (parallel)
• Hund's second rule states that the most stable configuration for an atom is the one with the maximum number of unpaired electrons allowed by the Pauli exclusion principle
• Hund's third rule states that in a set of degenerate orbitals, the configuration with the highest total spin multiplicity (2S+1) is the most stable

Applying Hund's rules

• To determine the ground state electron configuration, first fill the orbitals according to the Aufbau principle, then apply Hund's rules to distribute electrons among degenerate orbitals
• Hund's rules help predict the ground state electron configuration, which is the most stable and lowest energy state of an atom
• Example: The ground state electron configuration of nitrogen (atomic number 7) is 1s² 2s² 2p³, with three unpaired electrons in the 2p subshell, following Hund's rules

Atomic stability and reactivity

Stability based on electron configuration

• Atoms with completely filled valence shells (noble gas configurations) are generally more stable and less reactive than those with partially filled valence shells
• Atoms with half-filled or fully filled subshells (e.g., p³, d⁵, d¹⁰) have extra stability due to exchange energy and symmetry, making them less reactive
• Example: Neon (1s² 2s² 2p⁶) and argon (1s² 2s² 2p⁶ 3s² 3p⁶) are highly stable and unreactive due to their complete octet valence configurations

Reactivity based on electron configuration

• Atoms with one or two electrons more than a stable configuration (e.g., alkali and alkaline earth metals) are highly reactive due to their tendency to lose electrons and achieve a stable configuration
• Atoms with one or two electrons fewer than a stable configuration (e.g., halogens) are highly reactive due to their tendency to gain electrons and achieve a stable configuration
• The reactivity of an atom can be predicted by its electron configuration and its proximity to a stable (noble gas) configuration
• Example: Sodium (1s² 2s² 2p⁶ 3s¹) is highly reactive because it can easily lose one electron to achieve a stable neon-like configuration

Electron shielding and atomic properties

Concept of electron shielding

• Electron shielding, also known as screening, is the reduction of the effective nuclear charge (Zeff) experienced by an electron due to the repulsive forces of other electrons in the atom
• Inner shell electrons (core electrons) shield the outer shell electrons (valence electrons) from the full attractive force of the positively charged nucleus
• The shielding effect is more pronounced for electrons in the same principal quantum shell (n) and less pronounced for electrons in different shells

Effects on atomic properties

• Electron shielding affects the atomic radius, ionization energy, and electron affinity of an atom
• As the number of shielding electrons increases, the atomic radius increases, while the ionization energy and electron affinity decrease
• Electron shielding helps explain the periodic trends in atomic properties across the periodic table, such as the decrease in ionization energy and increase in atomic radius from top to bottom in a group
• Example: In the lithium atom (1s² 2s¹), the 2s electron is shielded by the two 1s electrons, resulting in a larger atomic radius and lower ionization energy compared to helium (1s²)