Electron configuration is a fundamental concept in chemistry that describes the arrangement of electrons in an atom. This arrangement determines how atoms interact with each other and explains much of the structure and logic of the periodic table. Understanding electron configuration helps predict chemical behavior, bonding tendencies, and element classification.

Atoms consist of a nucleus containing protons and neutrons, surrounded by electrons that occupy specific energy levels. These energy levels, or shells, are numbered starting from the one closest to the nucleus (n=1, n=2, etc.). Each shell can hold a certain number of electrons, determined by quantum mechanical principles.

Within each shell are subshells designated as s, p, d, and f. Each subshell consists of orbitals-regions in space where electrons are most likely to be found. The number of orbitals in a subshell determines how many electrons it can hold:
- s subshell: 1 orbital, 2 electrons
- p subshell: 3 orbitals, 6 electrons
- d subshell: 5 orbitals, 10 electrons
- f subshell: 7 orbitals, 14 electrons

Electrons fill orbitals in a specific order based on their energy. The Aufbau principle states that electrons occupy the lowest energy orbital available. Hund's rule and the Pauli exclusion principle also govern how electrons fill orbitals. Hund's rule states that electrons will fill degenerate orbitals (orbitals of the same energy) singly before pairing, and the Pauli exclusion principle states that no two electrons can have the same set of quantum numbers.

A typical electron configuration is written using the notation of energy level, subshell, and number of electrons. For example, the configuration of carbon (atomic number 6) is 1s^2 2s^2 2p^2. This indicates two electrons in the first shell's s orbital, and four electrons in the second shell, split between the s and p orbitals.

The periodic table is arranged in a way that reflects electron configurations. Elements are grouped into periods (rows) and groups (columns) based on similar electron configurations. For instance, all elements in Group 1 (alkali metals) have one electron in their outermost s orbital, making them highly reactive.

As you move across a period from left to right, electrons are added to the same shell, but the increasing number of protons pulls them closer to the nucleus, causing a decrease in atomic radius. As you move down a group, new energy levels are added, increasing the atomic radius.

Electron configuration also explains periodic trends:
- Ionization energy increases across a period and decreases down a group.
- Electronegativity follows a similar pattern.
- Atomic radius decreases across a period and increases down a group.

The noble gases at the end of each period have full outer electron shells, making them stable and largely unreactive. Elements that do not have full outer shells tend to gain, lose, or share electrons to achieve stability, following the octet rule.

Electron configurations are also essential for understanding chemical bonding. Atoms with similar valence electron configurations often form similar types of bonds. For example, halogens (Group 17) need one more electron to complete their outer shell and readily form compounds with alkali metals (Group 1), which have one electron to lose.

In summary, electron configuration is a key to understanding atomic structure, periodic table organization, and chemical reactivity. By mastering this concept, chemists can predict and explain the behavior of elements across the entire table.