<unk>  is an observational<unk>  of atomic chemistry discovered by Friedrich<unk> , which is one of a set of<unk> s referred to collectively as<unk> 's<unk> s. They are important for spectroscopy and quantum chemistry.<unk>  is a principle (1925) of atomic chemistry which states that a greater total spin state usually makes the resulting atom more stable, most commonly manifested in a lower energy state, because it forces the unpaired electrons to reside in different spatial orbitals. A commonly given reason for the increased stability of high<unk>  states is that the different occupied spatial orbitals create a larger average distance between electrons, reducing electron-electron repulsion energy. In reality, it has been shown that the actual reason behind the increased stability is a decrease in the screening of electron-nuclear attractions. Total spin state is calculated as the total number of unpaired electrons + 1, or twice the total spin + 1 written as 2s+1.
As a result of<unk> 's<unk> , constraints are placed on the way atomic orbitals are filled using the Aufbau principle. Before any two electrons occupy an orbital in a subshell, other orbitals in the same subshell must first each contain one electron.  Also, the electrons filling a subshell will have parallel spin before the shell starts filling up with the opposite spin electrons (after the first orbital gains a second electron). As a result, when filling up atomic orbitals, the<unk>  number of unpaired electrons (and hence<unk>  total spin state) is assured.
In 2004, researchers reported the synthesis of 5-dehydro-m-xylylene (DMX), the first organic molecule known to violate<unk> 's<unk> .
Hund's rule of maximum multiplicity
Cannon-Bard theory
Alignments of random points
Franck-Condon principle
Gauss–Markov theorem
Weber–Fechner law
Gibbs' phase rule
Ancillary statistic