Hund s rules

Hund s rules Rules which describe the electronic configuration of degenerate orbitals in the ground state. The electronic configuration will have the maximum number of unpaired... 

Procedure. To go from an STO-3G ealeulation to a CBS-4 ealeulation, simply replaee STO-3G with CBS-4 in the route seetion of the program used in Computer Projeet 8-1. Complete Table 8-2 by filling in the CBS-4 Energies of the atoms and ions listed in eolumns 1 and 3 of Table 8-2 and put them into eolumns 2 and 4 of the table. You will notiee that some of the simpler atoms (H through Be) do not have a listed CBS-4 Energies, but they do have an SCF energy, whieh should be used in its plaee. Caleulate the IP and eomplete eolumn 5. Pay speeial attention to spin niultiplieity and Hund s rule. The spin niultiplieity is rr + 1 where n is the number... 

All four sp orbitals are of equal energy Therefore according to Hund s rule (Sec tion 1 1) the four valence electrons of carbon are distributed equally among them making four half filled orbitals available for bonding... 

Hund s rule (Section 1 1) When two orbitals are of equal en ergy they are populated by electrons so that each is half filled before either one is doubly occupied Hybrid orbital (Section 2 6) An atomic orbital represented as a mixture of vanous contributions of that atom ss p d etc orbitals... 

Hund s rules tell us that the lowest energy term of these is as it has the highest spin multiplicity (25 +1=3) and the highest value of L (3 for an F term). 

Hund s rule (Section 1.3) If two or more empty orbitals of equal energy are available, one electron occupies each, with their spins parallel, until all are half-full. 

Humulene. structure of, 202 Hund s rule, 6 sp Hybrid orbitals. 17-18 sp2 Hybrid orbitals, 15. sp3 Hybrid orbitals, 12-14 Hydrate, 701... 

Similar situations arise frequently. There is a general principle that applies in all such cases Hund s rule (Friedrich Hund, 1896-1997) predicts that, ordinarily,... 

Hund s rule, like the Pauli exclusion principle, is based on experiment It is possible to determine the number of unpaired electrons in an atom. With solids, this is done by studying their behavior in a magnetic field. If there are unpaired electrons present the solid will be attracted into the field. Such a substance is said to be paramagnetic. If the atoms in the solid contain only paired electrons, it is slightly repelled by the field. Substances of this type are called diamagnetic. With gaseous atoms, the atomic spectrum can also be used to establish the presence and number of unpaired electrons. 

Strategy Start with the electron configuration, obtained as in Section 6.5. Then write the orbital diagram, recalling the number of orbitals per sublevel, putting two electrons of opposed spin in each orbital within a completed sublevel, and applying Hund s rule where sublevels are partially filled. 

The atomic number of iron is 26 its election configuration is ls22s22p63s23p64s23d6. All the orbitals are filled except those in the 3d sublevel, which is populated according to Hund s rule to give four unpaired electrons. 

Reality Check To construct an orbital diagram, start with the electron configuration and apply Hund s rule. 

Strategy First (1) find the total number of electrons (Z Co = 27). Then (2) find the electron configuration the first 18 electrons form the argon core, and the remaining electrons enter the 3d sublevel. Finally (3) apply Hund s rule to obtain the orbital diagram. 

When A0 is small, the electron distribution is the same as in the simple cation if A is large, Hund s rule is not strictly followed. 

The idealized configurations refer to those expected on the basis of Hund s rule, that is configurations in which spin multiplicity is maximized. The only elements where the idealized configurations are found to occur are Ce, Gd, and Lu. Many early assignments of rare earth configurations had assumed the above-given idealized versions, due to the predominant trivalency of the rare earths. This provides another example of a difference between the chemical and spectroscopic periodic tables. 

Electrons occupy orbitals in such a way as to minimize the total energy of an atom by maximizing attractions and minimizing repulsions in accord with the Pauli exclusion principle and Hund s rule. 

If more than one orbital in a subshell is available, add electrons to different orbitals of the subshell before doubly occupying any of them (Hund s rule). 

We account for the ground-state electron configuration of an atom by using the building-up principle in conjunction with Fig. 1.41, the Pauli exclusion principle, and Hund s rule. 

If more than one molecular orbital of the same energy is available, the electrons enter them singly and adopt parallel spins (Hund s rule). 

Hiroshima, 721 histidine, 443, 774 hole, 195 homeostasis, 386 HOMO, 126, 580 homogeneous alloy, 202 homogeneous catalyst, 565 homogeneous equilibria, 362 homogeneous mixture, F53 homolytic dissociation, 80 homonuclear diatomic molecule, 103 Hooke s law, 92 hormone, 670 horsepower, A4, 791 hour, A4 HPLC, 354 HRF products, 723 HTSC, 192 Humphreys series, 51 Hund, F 35 Hund s rule, 35, 37 Hurricane Rita, 144 hyaluronic acid, 344 hybrid orbital, 109 hybridization bond angle, 131 molecular shape, 111 hydrangea color, 463 hydrate, F32 hydrate isomer, 676 hydration, 178 hydrazine, 627... 

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]

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