Inorganic Chemistry: Electronic Configurations

Inorganic Chemistry: Electronic Configurations

Understanding Electronic Configurations

• The Electronic Configuration of an atom describes the distribution of electrons in its atomic orbitals.
• The atom’s energy levels, also known as shells, are filled starting with the lowest energy level.
• The arrangement of electrons follows the Aufbau principle that states that an electron occupies the lowest energy orbital available.
• Each shell is indicated by a specific quantum number (n), starting with n=1 for the innermost shell.

The Pauli Exclusion Principle

• The Pauli Exclusion Principle states no two electrons in an atom can have identical quantum numbers.
• In other words, in any given atomic orbital, a maximum of two electrons can exist and they must have opposite spin.
• Spin is also a quantum number and can have one of two values: +1/2 (spin up) or -1/2 (spin down).

Writing Electronic Configurations

• To write an electronic configuration, the shells (n) and subshells (s, p, d, f) are written with the number of electrons in each.
• For example, the electronic configuration of oxygen (atomic number 8) is 1s² 2s² 2p⁴.
• The superscript indicates the number of electrons in the particular subshell.

The after effects of Electron Configuration

• Electron configurations influence the behaviour of atoms in chemical reactions.
• The covalent bonding nature of elements is determined by the number and arrangement of their outer shell (or valence) electrons.
• The electronic configuration determines the atom’s chemical reactivity and the types of bonds it can form.

Exceptions to the Aufbau Principle

• There are certain exceptions to the Aufbau principle, which can be explained by the fact that half-filled and fully filled subshells are extra stable.
• Examples include chromium and copper, where an electron from a lower energy level is used to half-fill or fully fill a higher energy subshell to achieve extra stability.