Electrons Hund’s rule

When there is a total of four electrons, Hund s rule predicts that two will be in the lowest orbital but the other two will be unpaired, so that the system will exist as a diradical rather than as two pairs. The degeneracy can be removed if the molecule is distorted from maximum molecular symmetry to a structure of lesser symmetry. For example, if 44 assumes a rectangular rather than a square shape, one of the previously degenerate orbitals has a lower energy than the other and will be occupied by two electrons. In this case, of course, the double bonds are essentially separate and the molecule is still not aromatic. Distortions of symmetry can also occur when one or more carbons are replaced by hetero atoms or in other ways.124... 

The correct answer is (C). Paramagnetism can be seen in atoms with unpaired electrons. Hund s rule allows us to predict the pairing of electrons in atomic orbitals, and, therefore, the magnetic properties of the atom. 

Analyze and Plan Because oxygen has an atomic number of 8, each oxygen atom has 8 electrons. Figure 6.24 shows the ordering of orbitals. The electrons (represented as arrows) are placed in the orbitals (represented as boxes) beginning with the lowest-enei orbital, the Is. Each orbital can hold a maximum of two electrons (the Pauli exclusion principle). Because the 2p orbitals are degenerate, we place one electron in each of these orbitals (spin-up) before pairing any electrons (Hund s rule). 

Where there are states which differ only in the number of unpaired electrons, Hund s Rule states that the state with the maximum number of parallel electron spins has the lowest energy. 

It has been found that electrons behave as if they spin on an axis, and only electrons spinning in opposite directions (indicated by t and ) can occupy the same orbital. This principle, known as the Pauli exclusion principle, explains why orbitals can contain a maximum of two electrons. Hund s rule and the Pauli exclusion principle can be combined Electrons will pair with other electrons in an orbital only if there is no empty orbital of the same energy available and if there is one electron with opposite spin already in the orbital. 

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... 

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 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. 

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. 

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). 

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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