The arrangement of electrons within an atom dictates its chemical behavior. Understanding this structure requires electron configuration, which describes the placement of electrons into specific regions around the nucleus called orbitals. This organization follows established rules to achieve the most stable, lowest-energy state. Nitrogen, a fundamental element with an atomic number of 7, serves as an excellent case study for determining how many electrons reside in its \(2s\) orbital.
Defining Nitrogen’s Atomic Structure
A neutral nitrogen atom is defined by its atomic number, 7. This indicates the nucleus contains 7 protons, and the atom must possess 7 electrons to be electrically neutral. These 7 electrons must be positioned to minimize the atom’s overall energy, which is the ground state configuration.
Electrons occupy regions of probability described by quantum mechanics, not simple, fixed paths. These regions are grouped into principal energy levels, or shells (\(n=1, 2, 3,\) and so on). Each shell is divided into subshells, which contain orbitals. The goal of determining the electron configuration is to locate these 7 electrons within these available energy levels.
The Meaning of the 2s Orbital
The designation “\(2s\)” provides information about the orbital’s location and characteristics. The number “2” represents the principal quantum number, indicating that this orbital belongs to the second main energy level or electron shell. This shell is farther from the nucleus than the first shell (\(n=1\)).
The letter “\(s\)” denotes the angular momentum quantum number, specifying the orbital’s shape. An “\(s\)” orbital is spherically symmetric, meaning the probability of finding an electron is uniform in all directions from the nucleus. Crucially, any single orbital can hold a maximum of two electrons. This limit is imposed by the Pauli exclusion principle, requiring that electrons occupying the same orbital must have opposite spins.
Determining Nitrogen’s Electron Configuration
To find the number of \(2s\) electrons in nitrogen, the 7 total electrons must be systematically placed into the available orbitals starting from the lowest energy level upwards, following the Aufbau principle. The lowest energy orbital is the \(1s\) orbital, belonging to the first energy level.
The \(1s\) orbital can accommodate a maximum of two electrons, resulting in the notation \(1s^2\). After the lowest \(1s\) orbital is filled, the next available energy level is the \(2s\) orbital. The third and fourth electrons of nitrogen are then placed into this \(2s\) orbital.
Because the \(2s\) orbital is a single orbital, the Pauli exclusion principle allows it to hold exactly two electrons, both with opposite spins. Therefore, the configuration is built to \(1s^2 2s^2\), accounting for four of nitrogen’s seven electrons. The remaining three electrons are then placed into the next higher energy subshell, which is the \(2p\) subshell.
The complete electron configuration for nitrogen is \(1s^2 2s^2 2p^3\). This configuration clearly shows that the \(2s\) orbital is completely filled with two electrons.