Atomic orbitals, transition elements

The atoms of transition elements do not display the same general trend as the main group elements. A key reason for this is that electrons are added to inner energy levels — the d orbitals — rather than to the outer energy levels. As a result, Zgff changes relatively little, so atomic size remains fairly constant. In later chemistry courses, you will learn a more complete explanation for the atomic radii of transition-element atoms. 

The titanium atom has the electron configuration 3d2 4s2. Remembering that when atoms of transition elements lose electrons they are lost from the s orbital first, the Ti3+ ion has the configuration 3d1. The orbitals and electron populations for the Ti atom and Ti3+ ion can now be shown ... 

The size of atoms of transition elements within the same period changes only slightly because electrons are filling d orbitals rather than the ontermost energy level. Because the increase in nuclear charge is canceled by an increase in d electrons, the attraction of the valence electrons by the nuclens remains abont the same. Becanse there is little change in the nuclear attraction for the valence electrons, the atomic size remains relatively constant for the transition elements. 

X-ray Electromagnetic radiation of wave length c. 1 k. X-rays are generated in various ways, including the bombarding of solids with electrons, when they are emitted as a result of electron transitions in the inner orbits of the atoms bombarded. Each element has a characteristic X-ray spectrum. 

We will limit ourselves here to transition metals. It is well known that in these metals, the cohesive properties are largely dominated by the valence d electrons, and consequently, sp electrons can be neglected save for the elements with an almost empty or filled d valence shelP. Since the valence d atomic orbitals are rather localized, the d electronic states in the solid are well described in the tight-binding approximation. In this approximation, the cohesive energy of a bulk crystal is usually written as ... 

The values for the atomic saturation magnetization at the absolute zero, ferromagnetic metals iron, cobalt, and nickel are 2.22, 1.71, and 0.61 Bohr magnetons per atom, respectively.9 These numbers are the average numbers of unpaired electron spins in the metals (the approximation of the g factor to 2 found in gyromagnetic experiments shows that the orbital moment is nearly completely quenched, as in complex ions containing the transition elements). 

In this discussion of the transition elements we have considered only the orbitals (n— )d ns np. It seems probable that in some metals use is made also of the nd orbitals in bond formation. In gray tin, with the diamond structure, the four orbitals 5s5p3 are used with four outer electrons in the formation of tetrahedral bonds, the 4d shell being filled with ten electrons. The structure of white tin, in which each atom has six nearest neighbors (four at 3.016A and two at 3.17.5A), becomes reasonable if it is assumed that one of the 4d electrons is promoted to the 5d shell, and that six bonds are formed with use of the orbitals 4dSs5p35d. 

As the computational effort in the LDF approach grows, in the limit, only with the third power in the number of orbitals, it can be expected that fairly large systems with a hundred atoms, including transition metals, rare earth, and actinide elements, will become tractable. 

We note that the valence orbitals of metal atoms order in energy as AE>Ln>M. The d-levels of transition elements (M) range the lowest, and are therefore most sensitive for reduction, or to form a stable binary metal nitride. This may also explain the virtual absence of d-element compounds with 16 (valence) electron species, such as [N=N=N] , [N=C=N] , [N=B=N] T [C=C=CfT or [C=B=C] T at least through high-temperature syntheses. 

As the atomic number increases, so does the positive charge of the nucleus, and the electrons are bound with a higher energy. However, this increase is not linear. For example, the electrons in the d orbital of the third shell have a higher energy than those in the s orbital of the fourth shell, and hence the latter are filled first. The consequence is the unexpected behavior of the first ten transition elements. In the case of the actinides and lanthanides, even more inner orbitals are occupied. Nature is not so simple, but the scheme should help to visualize this complex structure. And if one can assign the electrons of an element, one is a step closer to successfully unraveling the mysteries of the Periodic Table. 

Thus, the calculations show that the outer ns(np) atomic orbitals can play a significant role in the formation of M-M bonds in transition metal acido-clusters. The probability that these atomic orbitals will participate in the formation of M-M bonds is maximal for elements of Group 7, particularly, for technetium, in whose clusters Zeff for technetium atoms is the lowest of those observed in all known acido-clusters. 

As can be seen from Table 1, not only the spectral data are quite different between pairs of compounds, but also the paramagnetism is decreasing when the carbon atom attached to the nitrogen is replaced by silicon, all other atoms being equal. As we have not been able to determine the molecular structures of the compounds until now, we cannot ascribe the change in properties to a definite change in structure. Nevertheless it seems obvious that the carbon or silicon atom in 6-position to the metal must have an important impact on the orbital-splitting at the transition element. 

Figure 4.93 illustrates some aspects of the break in the vertical trend of atomic orbital energies es and for early, middle, and late transition elements, showing the contrasting behavior of third-series versus first- and second-series elements. The... 

Figure Al.l Approximate energy level diagram for electronic orbitals in a multi-electron atom. Each horizontal line can accommodate two electrons (paired as so-called spin-up and spin-down electrons), giving the rules for filling the orbitals - two in the s-levels, 6 in the p-levels, 10 in the d-levels. Note that the 3d-orbital energy is lower than the 4p, giving rise to the d-block or transition elements. (From Brady, 1990 Figure 7.10. Copyright 1990 John Wiley Sons, Inc. Reprinted by permission of the publisher.)...

The principal characteristic of the transition elements is an incomplete electronic subshell that confers specific properties on the metal concerned. Ligand systems may participate in coordination not only by electron donation to the 3d levels in the first transition series but also by donation to incomplete outer 4s and 4p shells. Figure 5.1 shows that the differences in orbital energy levels between the 4s, 4p and 3d orbitals are much smaller than, for example, the difference between the inner 2s and 2p levels. Consequently, transitions between the 4s, 4p and 3d levels can easily take place and coordination is readily achieved. The manner in which ligand groups are oriented in surrounding the central metal atom is determined by the number and energy levels of the electrons in the incomplete subshells. 

The orbitals of the d states in clusters of the 3d, 4d, and 5d transition elements (or in the bulk metals) are fairly localized on the atoms as compared with the sp valence states of comparable energy. Consequently, the d states are not much perturbed by the cluster potential, and the d orbitals of one atom do not strongly overlap with the d orbitals of other atoms. Intraatomic d-d correlations tend to give a fixed integral number of d electrons in each atomic d-shell. However, the small interatomic d-d overlap terms and s-d hybridization induce intraatomic charge fluctuations in each d shell. In fact, a d orbital contribution to the conductivity of the metals and to the low temperature electronic specific heat is obtained only by starting with an extended description of the d electrons.7... 

For the representative elements, the valence electrons are all electrons in the outer s and p orbitals of an atom. A quick way of determining the number of valence electrons is to locate the element on the periodic table. There are eight columns of representative elements. The first column, headed by H and Li, has one valence electron, the second column has two, skip the transition elements, the next column, headed by B and Al, has three. This continues to the last (eighth) column where there are eight valence electrons. The only exception to this procedure is helium, which only has two valence electrons. 

The d block includes all the transition elements. In general, atoms of d block elements have filled ns orbitals, as well as filled or partially filled d orbitals. Generally, the ns orbitals fill before the (n - l)d orbitals. However, there are exceptions (such as chromium and copper) because these two sublevels are very close in energy, especially at higher values of n. Because the five d orbitals can hold a maximum of ten electrons, the d block spans ten groups. 

The /block includes all the inner transition elements. Atoms of /block elements have filled s orbitals in the outer energy levels, as well as filled or partially filled 4/and 5/orbitals. In general, the notation for the orbital filling sequence is ns, followed by (n - 2)/, followed by (n - l]d, followed by (for period 6 elements) np. However, there are many exceptions that make it difficult to predict electron configurations. Because there are seven/orbitals, with a maximum of fourteen electrons, the /block spans fourteen groups. 

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