Electron configuration notation illustrates how electrons are distributed among orbitals in an atom. This information helps in understanding chemical properties and reactivity.

The Aufbau Principle explains that electrons occupy lower-energy orbitals before filling higher-energy ones. This systematic filling is essential for determining the electron configuration of an atom.

The superscript shows how many electrons are present in a specific orbital. This detail is crucial for understanding the electron arrangement within an atom.

The notation '2s' indicates that the electron is located in the s orbital of the second energy level. This is part of the systematic approach to filling orbitals based on increasing energy levels.

A 1s orbital can hold a maximum of 2 electrons, which is represented as 1s². Any other configuration either underpopulates or overpopulates the orbital, violating the Pauli Exclusion Principle.

Sodium has 11 electrons, so following the Aufbau Principle the electrons fill in this order: 1s, 2s, 2p, then 3s. The configuration 1s² 2s² 2p❶ 3s¹ correctly sums to 11 electrons.

Hund's Rule states that electrons occupy degenerate orbitals singly with parallel spins to minimize repulsion. This arrangement lowers the energy and increases the stability of the atom.

The quantum numbers that describe an electron are the principal, azimuthal, magnetic, and spin quantum numbers. The mass number, however, relates to the atom as a whole and not to the electron's properties.

Within the Aufbau sequence, orbitals are filled in order of increasing energy. The 3p orbital is filled before the 4s orbital, making the 4s orbital higher in energy compared to 3p.

The Pauli Exclusion Principle mandates that every electron in an atom must have a unique set of quantum numbers. This rule is essential for the proper arrangement of electrons in orbitals.

The transition metals block is defined by the progressive filling of d orbitals. This unique filling pattern gives these elements their characteristic chemical and physical properties.

Each orbital, including a p orbital, can hold a maximum of 2 electrons with opposite spins due to the Pauli Exclusion Principle. Although a p subshell contains three orbitals (for a total of 6 electrons), each individual orbital is limited to 2 electrons.

For chlorine, which has 17 electrons, the electrons fill the orbitals in the order prescribed by the Aufbau Principle. The configuration 1s² 2s² 2p❶ 3s² 3p❵ correctly accounts for all 17 electrons.

A d subshell comprises 5 orbitals. Since each orbital can hold 2 electrons, a d subshell can accommodate a total of 10 electrons.

The electron spin quantum number can be either +½ or -½, indicating the intrinsic spin of an electron. This property influences magnetic behavior and is an essential factor in the Pauli Exclusion Principle.

Chromium exhibits an anomalous electron configuration to achieve extra stability with a half-filled d subshell. The configuration 1s² 2s² 2p❶ 3s² 3p❶ 4s¹ 3d❵ reflects this stability, differing from the expected order.

Copper has an anomalous electron configuration that maximizes stability by having a fully filled d subshell. The configuration 1s² 2s² 2p❶ 3s² 3p❶ 4s¹ 3d¹❰ is observed instead of the expected order.

For an electron in a 3d orbital, the principal quantum number (n) is 3 and the angular momentum quantum number (l) is 2. The magnetic quantum number (m) for a d orbital ranges from -2 to +2, making m = -1 valid along with a spin of +½.

Electrons fill orbitals in the order of increasing energy to achieve the lowest overall energy state for the atom. This systematic filling is described by the Aufbau principle, ensuring stability.

The electron configuration reveals the distribution of electrons, particularly those in the outer (valence) shell. These valence electrons are responsible for an element's chemical reactivity and bonding behavior.