The arrangement of electrons within an atom is known as its electron configuration. Understanding electron configurations is crucial in explaining the organization of elements in the periodic table.

The primary principles that govern electron configurations are the Aufbau principle, Hund's rule, and the Pauli exclusion principle. According to the Aufbau principle, electrons fill the lowest available energy levels before occupying higher energy levels. To illustrate this, let's consider the electron configuration of nitrogen (N), which has an atomic number of 7. The first and second energy levels, which are occupied by two electrons in total, are filled with the 1s^2 and 2s^2 orbitals. The remaining three electrons occupy the 2p orbital, with one electron in each of the 2p_x, 2p_y, and 2p_z orbitals.

Hund's rule states that when filling orbitals with degenerate energies (same energy level), electrons prefer to occupy separate orbitals with parallel spins rather than pairing up. This is evident in the 2p orbital of nitrogen, where each electron has an unpaired spin.

The Pauli exclusion principle states that no two electrons in an atom can have the same four quantum numbers (n, l, m, and s). This means that each electron in an atom has a unique set of quantum numbers that determines its location and spin.

By understanding and applying these principles, we can determine the electron configurations of different elements and use this information to explain their positions in the periodic table.