B.sc 1st year Book
(Page 8)

# Hunds Rule of Maximum Multiplicity

On the basis of magnetic measurements, Hund postulated an important empirical rule popularly known after his name as Hund’s Rule. This rule has a spectroscopic basis and is applied when orbitals of a sub-group of an atom are incompletely occupied. This rule may be defined as :

Electrons prefer to occupy separate orbitals so that they have parallel spins:

or’ When electrons are added in a set of orbitals in a given subshell (p- d- or f-) never pair up until all the available orbitals do not have one electron each.
(As students in a classroom prefer to occupy a single separate seat until all the chairs do not fill). Thus, it is a natural law.

Both the above statements have the same meaning that the pairing of electrons will occur in a set of orbitals of equivalent energies only when all the available orbitals do not contain one electron each. The atoms, ions, or molecules obeying the above statements are known to be thermodynamically more stable.

### Hunds Rule Explanation :

To explain Hund’s rule, let us consider there are three electrons that are allowed to enter p-orbitals, their distribution will be:

Thus, it is clear from the above configurations that pairing begins with the introduction of the 2nd electron in the s-orbital, the 4th electron in p-orbitals, the 6th in d-orbitals, and the 8th electron in f-orbitals. These can be shown by the following representations.

Hund’s rule also states that an atom containing empty  (p°, d°, or f°), completely half-filled, or completely full-filled (p3, d5, or f7) orbitals will have more stability as compared to other arrangements. Because such arrangements provide symmetry. In other words, empty, completely half-filled, and completely full-filled orbital provide extra stability to an atom or other species.

## Related Topic | Atomic Structure and Periodic Table

 Bohr’s Atomic Model De Broglie Hypothesis De Broglie Equation Heisenberg uncertainty Principle Schrodinger Wave Equation Shapes of orbitals Quantum Numbers Hunds Rule Pauli Exclusion Principle Energy level diagram of Atom Aufbau Principle Screening effect or shielding effect Periodic Table Moseley Periodic law Classification of elements

### Hunds Rule Importance :

(i) Hund’s rule plays an important role in interpreting the diamagnetic and paramagnetic properties of various ions, molecules, and complex compounds. For example, a Cupprus ion Cu++ consisting of no unpaired electron is diamagnetic whereas a cuprous ion Cu++ is paramagnetic due to the presence of one unpaired d-electron as clear from the following electronic configurations:

(ii) This rule satisfactorily explains how an atom exhibits more than one electro-valency for metals and co-valency for non-metals. For example :

(a) Copper exhibits +1 electro-valency when it follows Hund’s rule and +2 when it does not follow Hund’s rule.
(b) Phosphorus has minimum covalency of 3 in-ground states and a maximum covalency of 5 in an excited state.

(iii) This rule explains the relative stability of the Oxyacids of transition metals.
(iv) Hund’s rule also explains the possible number of states or levels of various species.

The Hunds rule develop. When scientists hunds working on electron distribution in orbitals. He finds that each orbital has a different energy level occupied by different no. of electrons on the different shells. Where he observed that each subgroup of orbital 1st occupied a single electron on every subgroup of the orbital. Then it’s occupying double.

 ↑↓ ↑↓ ↑↓ ↑ 1s 2s 2px 2py 2pz X(Wrong) ↑↓ ↑↓ ↑ ↑ ↑ 1s 2s 2px 2py 2pz √(Right)

Hund’s rule: Each subshell of orbitals is first singly occupied or filled. Then, it’s going to occupy another half of the subshell. This means that the first half subshell will be filled and then doubled.

State of Hund’s Rule :

1. The orbital of the subgroup is occupied signal 1st in every subshell or subgroup. Then it goes doubly occupied.

2. The orbital of a subshell or subgroup which is singly occupied or half filled have their identical skin.

Example 1: Fluorine (F)

We can consider the correct electron configuration of the fluorine (Z=9) Atom: 1s2 2s2 2p5

 Fluorine
 ↑↓ ↑↓ ↑↓ ↑↓ ↑ 1s2 2s2 2p5

The p orbitals of the subgroup or subshell are half-filled; there are 5 electrons and three p orbitals. This is because the 5 electrons in the 2p subshell/subgroup will fill all the empty orbitals 1st before pairing with e electrons in them.

Reminder in mind that elemental fluorine is found in nature typically or hard to find as molecular fluorine, F2, which has molecular orbitals instead of atomic orbitals as demonstrated above.

Example 2: Sodium (Na)

We can consider the correct electron configuration of the fluorine (Z=11) Atom: 1s2 2s2 2p6 3s1

 Sodium
 ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ 1s2 2s2 2p6 3s1

The p orbitals of the subgroup or subshell are filled; there are 5 electrons and three p orbitals and 1 electron in 3s orbitals which is half filled. This is because the 1 electron in the 3s subshell/subgroup will fill all the empty orbitals 1st before pairing with e- electrons in them.

Reminder in mind that elemental fluorine is found in nature typically or hard to find as molecular Sodium, (Na)2, which has molecular orbitals instead of atomic orbitals which is hard to find it shown as demonstrated above