Hund’s rule explains electron arrangement in an atom’s orbitals. Each orbital must contain one electron before pairing, and unpaired electrons must spin in the same direction. This rule is important for understanding magnetism in metals.
Used in physics, Hund’s rule concerns 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 the orbitals. The rule is important for understanding certain behaviors in atoms, such as magnetism in 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 are negatively charged. Electrons move, or spin, in certain areas around the nucleus, called orbitals, and can have another electron share their orbit. When this happens, the electrons will spin in opposite directions.
Besides spins, electron orbitals are also defined by subshells and energy levels. Subshells are labeled with the letters s, p, d, and denote certain orbitals or groups of orbitals that 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 contains only a subshell s, the second energy level has a subshell s and a subshell ap, and so on. Simply put, the more electrons an atom has, the more subshells and energy levels it has.
For example, hydrogen contains only one electron, so it has only one subshell, the s, in the first energy level. In contrast, iron contains 26 electrons, so it has four subshells, one for each energy level; two p subshells, each containing three orbitals, located in energy levels two and three; and a d subshell, containing five orbitals, in energy level three.
Focusing on the outer shell, Hund’s rule determines how electrons are arranged in orbitals or their configuration. Based on 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. He also says that as empty orbitals fill up, each unpaired electron must spin in the same direction. Since a subshell must be completely full before electrons fill other shells, this rule only applies to the last filled subshell.
For example, iron’s 26 electrons completely fill each of its subshells up to the last, the 3d subshell. Here are six electrons left to fill five orbitals. The first five electrons, all spinning in the same direction, will each occupy one orbital, and the sixth will pair with the electron in the first orbital, spinning in the opposite direction. It is this phenomenon, with an unpaired number of electrons all spinning in the same direction, that allows objects to become magnetic. Conversely, when all the electrons in the outer shell are paired, as with the noble gases, the atoms are completely stable.
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