Electronic Configuration

• The electronic configuration of an atom represents the arrangement of electrons within it. Electrons are situated in shells, or energy levels, orbiting the nucleus of the atom.
• There are rules that govern the ordering of electrons, known as the aufbau principle. According to this principle, electrons occupy the orbitals of lowest energy first.
• Each shell can hold a fixed number of electrons determined by the 2n^2 principle: the first shell (n=1) can hold up to 2 electrons, the second shell (n=2) can hold up to 8 electrons, and so on.
• An orbital is a region of space within an energy shell where there is a high probability of finding an electron. Each orbital can accommodate a maximum of two electrons.
• Orbitals are classified into types known as s, p, d, and f. The s orbital is spherically-shaped and can hold up to 2 electrons. The p orbital is dumbbell-shaped and can hold up to 6 electrons. The d orbital is cloverleaf-shaped and can hold up to 10 electrons. The f orbital is complex in shape and can hold up to 14 electrons.
• The Pauli Exclusion Principle states that in one orbital, there can be at most two electrons, and they must be spinning opposite to each other (one up-spin electron and one down-spin electron).
• Electronic configurations are often written using noble gas shorthand, where the previous noble gas in the Periodic Table stands for the electrons in the filled energy levels.
• Knowledge of electronic configuration helps understand chemical behaviour. Atoms with full outer energy levels are often unreactive, while those with space for more electrons in their outer levels tend to be more reactive.
• Chemical properties of elements are largely determined by the number of electrons in their outermost shell, also referred to as ‘valency’.
• Periods in the Periodic Table represent energy levels; each period corresponds to a new energy level being filled with electrons. Groups in Periodic Table represent the number of electrons present in the outermost shell.
• Electron configurations illustrate the principles of periodicity and Periodic Table trends: elements in the same group have the same number of valence electrons, which are responsible for chemical reactivity and similarity within a group.
• Transition metals will often lose electrons from the 4s orbital before the 3d, even though 4s is filled before 3d. This is related to the subtle differences in energy levels between these orbitals.

Remember, the more you practise writing electron configurations, the more confidence you will gain in this crucial skill for understanding atomic structure and chemical behaviour!