An atom’s chemical behavior depends on the spatial arrangement of its electrons. Electrons reside in complex, three-dimensional probability zones called orbitals, not simple planetary paths. Scientists use an orbital diagram to visually map the location and directional spin of electrons. Nitrogen (N) has an atomic number of seven, meaning a neutral atom possesses seven electrons. Determining the lowest-energy, or ground state, configuration for these electrons requires a systematic process guided by quantum mechanics.
Understanding Atomic Orbitals and Subshells
Electrons occupy specific regions of space around the nucleus grouped into distinct energy levels, or shells. Within these shells are subshells, which contain one or more orbitals. The types of orbitals are categorized by shape and designated by the letters s, p, d, and f. Since nitrogen is a small atom, only the s and p subshells are necessary to accommodate its seven electrons.
The s-orbital is the simplest shape, resembling a perfect sphere centered on the nucleus. Every energy level contains one s-orbital, which can hold a maximum of two electrons. The next highest energy subshell is the p-subshell, consisting of three individual p-orbitals.
These three p-orbitals are identical in size and energy, but they are oriented along the three perpendicular axes in space, often labeled \(p_x\), \(p_y\), and \(p_z\). Each of these three dumbbell-shaped p-orbitals can hold two electrons, giving the entire p-subshell a total capacity of six electrons. The orbital diagram visually represents these containers, using lines or boxes to denote each individual orbital within a subshell.
The Three Key Rules Governing Electron Placement
The systematic filling of atomic orbitals is governed by three principles that ensure the electron arrangement represents the atom’s most stable, lowest-energy state. The first is the Aufbau principle, which dictates the sequence for filling energy levels. This rule requires that electrons occupy the lowest-energy orbitals available before moving into higher-energy orbitals. For nitrogen, this means the \(1s\) orbital is filled first, followed by the \(2s\) orbital, and finally the \(2p\) orbitals.
The second principle is the Pauli exclusion principle, which addresses how many electrons can occupy a single orbital and their spin state. This means that any single orbital can contain a maximum of two electrons, and those two electrons must have opposite spins. This difference in spin is commonly represented by arrows pointing in opposite directions (up and down) within the orbital box.
The final rule is Hund’s Rule, which applies to subshells that contain multiple degenerate orbitals, such as the three \(2p\) orbitals. It states that when electrons are added to these equal-energy orbitals, they must first occupy each orbital singly. Furthermore, all unpaired electrons must have parallel spins, meaning their arrows point in the same direction. This arrangement minimizes repulsion between electrons, resulting in a more stable, lower-energy state for the atom.
Constructing the Orbital Diagram for Nitrogen
Applying the rules to the seven electrons of a neutral nitrogen atom begins by following the established energy sequence. The first two electrons enter the \(1s\) orbital, the lowest energy level. According to the Pauli exclusion principle, these two electrons must be paired and possess opposite spins, fully filling the \(1s\) orbital.
The next two electrons move to the \(2s\) orbital. The \(2s\) orbital is filled with a pair of electrons of opposite spin. At this stage, four of the seven electrons are accounted for, leaving three electrons for the three degenerate \(2p\) orbitals.
The last three electrons are placed into the \(2p\) subshell, where Hund’s rule becomes relevant. The rule requires that one electron is placed into each of the three \(2p\) orbitals. To achieve the lowest energy configuration, all three unpaired electrons must have the same spin direction, typically represented as three parallel up arrows. The resulting ground state orbital diagram for nitrogen shows \(1s\) and \(2s\) paired, and \(2p\) with three separate, singly-occupied orbitals with parallel spin.