What is a hunds rule?

Answer 1

Hund's rule of maximum multiplicity describes the order in which electrons fill subshells. It states that they will add into degenerate(equal energy level) orbitals to create the greatest multitude of orbitals having an electron in them.

Simply, this means that if you were filling the 2p subshell you would first put an electron in the 2px then 2py then 2pz THEN a second electron in 2px then a second in 2py then one in 2pz NOT 1 in 2px 2nd in 2px etc.

The orbitals of a subshell must be occupied singly and with parallel spins before they occupy in pairs.

Hund's rule states that 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.

Answer 2

When there there are degenerate (equal energy) orbitals available, electrons will enter the orbitals one-at-a-time to maximise degeneracy, and only when all the orbitals are half filled will pairing-up occur. This is the rule of maximum multiplicity.

Answer 3

Hund's Rule of Maximum Multiplicity is an observational rule which states that a greater total spin state usually makes the resulting atom more stable. Accordingly, it can be taken that if two or more orbitals of equal energy are available, electrons will occupy them singly before filling them in pairs.

The rule is used while filling of electrons in different orbitals.

Answer 4

In atomic physics, Hund's rules refer to a set of rules formulated by German physicist Friedrich Hund around 1927, which are used to determine the term symbol that corresponds to the ground state of a multi-electron atom. In chemistry, the first rule is especially important and is often referred to as simply Hund's Rule.

The three rules are: [1][2][3]

1. For a given electron configuration, the term with maximum multiplicity has the lowest energy. The multiplicity is equal to , where is the total spin angular momentum for all electrons. The term with lowest energy is also the term with maximum .
2. For a given multiplicity, the term with the largest value of the orbital angular momentum number has the lowest energy.
3. For a given term, in an atom with outermost subshell half-filled or less, the level with the lowest value of the total angular momentum quantum number (for the operator ) lies lowest in energy. If the outermost shell is more than half-filled, the level with the highest value of is lowest in energy.

These rules specify in a simple way how the usual energy interactions dictate the ground state term. The rules assume that the repulsion between the outer electrons is very much greater than the spin-orbit interaction which is in turn stronger than any other remaining interactions. This is referred to as the LS coupling regime.

Full shells and subshells do not contribute to the quantum numbers for total S, the total spin angular momentum and for L, the total orbital angular momentum. It can be shown that for full orbitals and suborbitals both the residual electrostatic term (repulsion between electrons) and the spin-orbit interaction can only shift all the energy levels together. Thus when determining the ordering of energy levels in general only the outer valence electrons need to be considered.

Rule 1

Due to the Pauli exclusion principle, two electrons cannot share the same set of quantum numbers within the same system; therefore, there is room for only two electrons in each spatial orbital. One of these electrons must have (for some chosen direction z) Sz = ½, and the other must have Sz = −½. Hund's first rule states that the lowest energy atomic state is the one which maximizes the sum of the S values for all of the electrons in the open subshell. The orbitals of the subshell are each occupied singly with electrons of parallel spin before double occupation occurs. (This is occasionally called the "bus seat rule" since it is analogous to the behaviour of bus passengers who tend to occupy all double seats singly before double occupation occurs.)

Two physical explanations have been given[4] for the increased stability of high multiplicity states. In the early days of quantum mechanics, it was proposed that electrons in different orbitals are further apart, so that electron-electron repulsion energy is reduced. Accurate quantum-mechanical calculations (starting in the 1970s) have shown that the reason is that the electrons in singly occupied orbitals are less effectively screened or shielded from the nucleus, so that such orbitals contract and electron-nucleus attraction energy becomes greater in magnitude (or decreases algebraically).

example: As an example, consider the ground state of silicon. The electronic configuration of Si is 1s2 2s2 2p6 3s2 3p2 (see spectroscopic notation). We need consider only the outer 3p2 electrons, for which it can be shown (see term symbols) that the possible terms allowed by the Pauli exclusion principle are 1D , 3P , and 1S. Hund's first rule now states that the ground state term is 3P which has S = 1. The superscript 3 is the value of the multiplicity = 2S + 1 = 3. The diagram shows the state of this term with ML = 1 and MS = 1.

Rule 2This rule deals with reducing the repulsion between electrons. It can be understood from the classical picture that if all electrons are orbiting in the same direction (higher orbital angular momentum) they meet less often than if some of them orbit in opposite directions. In the latter case the repulsive force increases, which separates electrons. This adds potential energy to them, so their energy level is higher. ExampleFor silicon there is no choice of triplet states , so the second rule is not required. The lightest atom which requires the second rule to determine the ground state is titanium (Ti, Z = 22) with electron configuration 1s2 2s2 2p6 3s2 3p6 4s2 3d2. Following the same method as for Si, the allowed terms include three singlets (1S, 1D, and 1G) and two triplets (3P and 3F). We deduce from Hund's first rule that the ground state is one of the two triplets, and from Hund's second rule that the ground state is 3F (with ) rather than 3P (with ). There is no 3G term since its state would require two electrons each with , in violation of the Pauli principle

Answer 5

The increased stability of the atom, most commonly manifested in a lower energy state, arises because the high-spin state forces the unpaired electrons to reside in different spatial orbitals. A false, but commonly given, reason for the increased stability of high multiplicity 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.[2] 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 Hund's rule, 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 maximum number of unpaired electrons (and hence maximum total spin state) is assured.

For example a p4 subshell arranges its electrons as [↑↓][↑][↑] rather than [↑↓][↑][↓] or [↑↓][↑↓][ ].

Answer 6

Hund's rule states that 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.

Answer 7

Hund's rule states that electrons will fill orbitals of a subshell singly before pairing up, and all orbitals of a given subshell will be half-filled before any are completely filled. This minimizes electron-electron repulsion and leads to greater stability in the atom.