Hund second rule

In Fig. 1 there is indicated the division of the nine outer orbitals into these two classes. It is assumed that electrons occupying orbitals of the first class (weak interatomic interactions) in an atom tend to remain unpaired (Hund s rule of maximum multiplicity), and that electrons occupying orbitals of the second class pair with similar electrons of adjacent atoms. Let us call these orbitals atomic orbitals and bond orbitals, respectively. In copper all of the atomic orbitals are occupied by pairs. In nickel, with ou = 0.61, there are 0.61 unpaired electrons in atomic orbitals, and in cobalt 1.71. (The deviation from unity of the difference between the values for cobalt and nickel may be the result of experimental error in the cobalt value, which is uncertain because of the magnetic hardness of this element.) This indicates that the energy diagram of Fig. 1 does not change very much from metal to metal. Substantiation of this is provided by the values of cra for copper-nickel alloys,12 which decrease linearly with mole fraction of copper from mole fraction 0.6 of copper, and by the related values for zinc-nickel and other alloys.13 The value a a = 2.61 would accordingly be expected for iron, if there were 2.61 or more d orbitals in the atomic orbital class. We conclude from the observed value [Pg.347]

Hund s second rule If ambiguity remains, the maximum-spin ground term will then be one with maximum L... 

The second step is called Hund s rule, for the German spectroscopist Friedrich Hund, who first proposed it. This procedure gives the configuration of the atom that corresponds to the lowest total energy, taking into... 

If two or more degenerate orbitals are available, one electron goes into each until all are half-full, a statement called Hund s rule. Only then does a second electron fill one of the orbitals. Furthermore, the electrons in each of the singly occupied orbitals must have the same value for their spin quantum number. 

The filling of the 3d subshell generally proceeds according to Hund s rule (Section 5.12) with one electron adding to each of the five 3d orbitals before a second electron adds to any one of them. There are just two exceptions to the expected regular filling pattern, chromium and copper ... 

Closely related to the Pauli exclusion principle is the third rule, Hund s rule, which states that when electrons occupy orbitals of equal energy (e.g., the five 3d orbitals), one electron enters each orbital until all the orbitals contain one electron. In this configuration, all electrons will have parallel spin (same direction). Second electrons then add to each orbital so that their spins are opposite to the first electrons in the orbital. Atoms with all outer orbitals half-filled are very stable. 

The correct answer is (B). The second electron in the p orbital should go into a different orbital. According to Hund s rule, all orbitals should fill with a single electron (each having like spins) prior to a second electron entering any orbital. 

Notice in Table 1-1 that carbon s third and fourth valence electrons are not paired they occupy separate orbitals. Although the Pauli exclusion principle says that two electrons can occupy the same orbital, the electrons repel each other, and pairing requires additional energy. Hund s rule states that when there are two or more orbitals of the same energy, electrons will go into different orbitals rather than pair up in the same orbital. The first 2p electron (boron) goes into one 2p orbital, the second 2p electron (carbon) goes into a different orbital, and the third 2p electron (nitrogen) occupies the last 2p orbital. The fourth, fifth, and sixth 2p electrons must pair up with the first three electrons. 

When occupying two or more orbitals with the same energy (for example, any of the three 2p orbitals), electrons will half fill each orbital until all are half filled before adding a second electron to each one. This is called Hund s rule. 

One of the early great triumphs of atomic theory was the aufbau principle which explained the periodic table. In it the atomic orbital energies were assigned the following order Is < 2s < 2p < 3s <. ..The occupation number of freeon orbitals was limited to no more than two as was discussed in Section 1. Further Hund s rule was imposed i. e. states of highest spin lie the lowest. The Gel fand state reconstruction of the aufbau for the second row of the periodic table (with mi = +1 or zero) is shown in Fig. 10.1 where 2p+ = + 1,2po = 0 and 2p.i = -1 ... 

According to Hund s rule, the electrons fill their orbitals singly and then they begin to pair up. This means that the second situation shown above is the correct filling order for the electrons. 

L takes the maximum value consistent with Hund s second rule For states of the same spin multiplicity, the state with the greater orbital angular momentum will... 

The Nd has the electron configuration [Xe]4f. Because it is a rare-earth element, spin-orbit coupling would be expected and hence, Eqs. 8.24-8.25 to apply. Furthermore, crystal-field splitting is usually unimportant for rare-earth ions because their partially filled 4f shells lie deep inside the ions, beneath filled 5s and 5p shells. Thus, the seven f orbitals would be degenerate and their occupancy would be a high-spin configuration, with the maximum value of S and L, in accordance with Hund s first and second rules ... 

The Russell-Saunders states that arise from interelectronic interactions in a free ion with an electronic configuration d are F, D, P, G, and S, with the F being the ground state as dictated by Hund s first and second rules. An approximate calculation in which the F state is outlined here. 

The energies of the various Russell-Saunders states of an atom or ion are not the same, and Hund s first and second rules can be used to decide which is the ground state ... 

Does an electron enter the first 3p orbital to pair with a single electron that is already there Or does the electron fill another 3p orbital According to Hund s rule, the second answer is correct. Hund s rule states that orbitals of the same n and I quantum numbers are each occupied by one electron before any pairing occurs. For example, sulfur s configuration is shown by the orbital diagram below. Electrons are represented by arrows. Note that an electron fills another orbital before the electron occupies an orbital that occupied. 

RULE 3 If two or more orbitals of equal energy are available, one electron occupies each until all orbitals are half-full. Only then does a second electron occupy one of the orbitals Hund s rule). The electrons in the half-filled orbitals all have the same spin. 

Hund s rule The fact that negatively charged electrons repel each other has an important impact on the distribution of electrons in equal-energy orbitals. Hund s rule states that single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same orbitals. For example, let the boxes below represent the 2p orbitals. One electron enters each of the three 2p orbitals before a second electron enters any of the orbitals. The sequence in which six electrons occupy three p orbitals is shown below. 

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