Give the ellectron arrangement of the periodic table

Electrons occupy discrete energy levels, or shells, denoted by the principal quantum number $n$ (n=1, 2, 3, ...). Higher $n$ values correspond to higher energy levels and shells further from the nucleus.

Within each shell, there are subshells (s, p, d, f) which correspond to different orbital shapes and energies.

• $n=1$: 1s (2 electrons)
• $n=2$: 2s (2 electrons), 2p (6 electrons)
• $n=3$: 3s (2 electrons), 3p (6 electrons), 3d (10 electrons)
• $n=4$: 4s (2 electrons), 4p (6 electrons), 4d (10 electrons), 4f (14 electrons) Each orbital can hold a maximum of two electrons with opposite spins (Pauli Exclusion Principle).

Electrons generally fill orbitals in order of increasing energy, starting from the lowest energy level. This order is approximately 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.

For degenerate orbitals (orbitals of the same energy, e.g., the three p orbitals), electrons will occupy each orbital singly with parallel spins before any orbital is doubly occupied.

The periodic table is structured based on the subshell being filled by the outermost (valence) electrons:

• s-block: Groups 1 and 2 (alkali metals and alkaline earth metals) where the outermost electrons fill an s orbital.
• p-block: Groups 13 to 18 (excluding helium) where the outermost electrons fill p orbitals.
• d-block: Groups 3 to 12 (transition metals) where the outermost electrons fill d orbitals.
• f-block: Lanthanides and actinides (inner transition metals) where the outermost electrons fill f orbitals.

• Periods (rows): Indicate the principal quantum number ($n$) of the highest occupied energy shell. For example, elements in Period 3 have their valence electrons in the $n=3$ shell.
• Groups (columns): Elements in the same group have similar valence electron configurations, leading to similar chemical properties. The group number often corresponds to the number of valence electrons (for main group elements).