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Used in physics, Hund’s rule deals with the arrangement of electrons in the orbitals of an atom. Hund’s rule states 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 into orbitals. The rule is important for understanding certain behaviors of atoms, such as the magnetism of metals.
At 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 have a negative charge. Electrons move, or rotate, in certain areas around the nucleus, called orbitals, and may have another electron sharing their orbit. When this happens, the 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, and f and denote certain orbitals or groups of orbitals that occur at different energy levels of atoms. There are four levels of ground state energy, which contain more sublayers as they increase. For example, the first energy level contains only an s sublayer, the second energy level has a s sublayer and an ap sublayer, and so on. Simply put, the more electrons an atom has, the more subshells and energy levels present.
For example, hydrogen contains only one electron, so it has only one subshell, os, in the first energy level. Iron, on the other hand, contains 26 electrons, so it has four subshells, one for each energy level; two p subshells, each containing three orbitals, located at energy levels two and three; and a d subshell, containing five orbitals, at energy level three.
Focusing on the outer shell, Hund’s rule determines how electrons are arranged in orbitals, or their configuration. Starting from the concept 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 all empty orbitals in a subshell before pairing with electrons. It also says that when filling the empty orbitals, each unpaired electron must rotate in the same direction. Since a subshell must be completely filled before electrons fill other shells, this rule only takes effect on the last filled subshell.
For example, iron’s 26 electrons fill each of its subshells completely down to the last one, the 3d subshell. Here, six electrons are left to fill five orbitals. The first five electrons, all spinning in the same direction, will each occupy an orbital, and the sixth will pair with the electron in the first orbital, spinning in the opposite direction. It is this phenomenon, with multiple unpaired electrons spinning in the same direction, that allows items to become magnetic. On the other hand, when all the electrons in the outer shell are paired, as with noble gases, the atoms are completely stable.