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Used in physics, Hund's rule addresses the arrangement of electrons in the orbitals of an atom. Hund's rule indicates that for any group of orbitals, or subshells, in an energy level, each orbital must contain one electron, each spinning in the same direction, before electrons can be paired in the orbitals. The rule is important to understanding certain behaviors in atoms, such as magnetism in metals.
In the center of an atom is the nucleus. The nucleus contains particles called protons — which are positively charged — and particles called neutrons, which are neutral. Moving around the nucleus are tiny particles called electrons, which are negatively charged. Electrons move, or spin, in given areas around the nucleus, called orbitals, and they may have one other electron sharing its orbit. When this happens, electrons will spin in opposite directions.
In addition to spins, electron orbitals are also defined by subshells and energy levels. Subshells are labeled with the letters s, p, d, and f, and denote certain orbitals or groups of orbitals which occur in the different energy levels of atoms. There are four ground state energy levels, which contain more subshells as they increase. For example, the first energy level only contains an s subshell, the second energy level has an s subshell and a p subshell, and so on. Put simply, the more electrons an atom possesses, the more subshells and energy levels present.
For example, hydrogen only contains one electron, so it only has one subshell, the s, in the first energy level. Conversely, iron contains 26 electrons, so it has four s subshells, one for each energy level; two p subshells, which each contain three orbitals, located in energy levels two and three; and one d subshell, containing five orbitals, in energy level three.
Focusing on the outer shell, Hund's rule determines how the electrons are arranged in the orbitals, or their configuration. Building off of the concepts that only two electrons can occupy a given orbital and electrons in the same orbital spin in opposite directions, Hund's rule states that electrons must always fill up all the empty orbitals in a subshell before pairing with electrons. It also says that when filling up the empty orbitals, every unpaired electron must spin the same direction. Since a subshell must be completely full before electrons fill other shells, this rule only really comes into effect for the last subshell filled.
For example, iron's 26 electrons fill each of its subshells fully until the last, the 3d subshell. Here, there are six electrons left to fill five orbitals. The first five electrons, all spinning the same direction, will each occupy an orbital, and the sixth will pair with the electron in the first orbital, spinning the opposite direction. It is this phenomenon, with a number of unpaired electrons all spinning the same direction, that allows items to become magnetic. Conversely, when all the electrons in the outer shell are paired, such as with the noble gases, the atoms are completely stable.