Transition Metal

What is a Transition Element? The transition elements comprise groups 3 to 12 and are found in the central region of the standard periodic table, an example of which is reproduced on the endpaper. This group is further subdivided into those of the first row (the elements scandium to zinc), the second row (the elements yttrium to cadmium) and the third row (the elements lanthanum to mercury). The term 'transition' arises from the elements' supposed transitional positions between the metallic elements of groups 1 and 2 and the predominantly non-metallic elements of groups 13 to 18. Nevertheless, the transition elements are also, and interchangeably, known as the transition metals in view of their typical metallic properties. The chemistry of the transition elements has been investigated for two centuries, and in the past fifty years these elements and their compounds have proved to be a nearly ideal touchstone for many of the models which have been developed to understand structure and bonding. The elements range from the widespread to the extremely rare; iron is the fourth most abundant element (by weight) in the earth's crust, technetium does not occur naturally. Elements such as gold and silver have been known in the native state since antiquity, whereas technetium was first prepared in 1937. Most of the elements exhibit a typical silvery metallic appearance, but gold and copper are unique in their reddish coloration and mercury is the only metal which is liquid at ambient temperatures. Compounds of the transition elements account for the majority of coloured inorganic materials, and many pigments are relatively simple derivatives of these elements; however, not all transition-element compounds are coloured. What are the common features that unite these elements? It is surprisingly difficult to find a single definition which satisfactorily encompasses all of the transition elements. The elements occur at that point in the periodic table where the d orbitals are being filled. The first row transition elements coincide with the filling of the 3d, the second row with the filling of the 4d, and the third row with the filling of the 5d orbitals. We define a transition element as possessing filled or partially filled valence d orbitals in one or more of its oxidation states. This definition excludes the elements in groups 13 to 18. The electron configurations of the transition elements are presented in Table 1-1. The outer configurations of the transition metals in Table 1-1 imply, and detailed spectroscopic investigations confirm, that the 3d orbitals lie at higher energies than the 4s orbitals.reveal the loss of electrons from the 4s shell in preference to the 3d, so that in these species the 4s orbitals are the higher in energy. The explanation of these facts is not difficult but is subtle. We recall that the energies of all hydrogen orbitals belonging to the same principal quantum shell (n) are equal: the 3d, 3p and 3s hydrogen orbitals are degenerate. lose their degeneracy, however, in many-electron atoms. Orbitals with smaller orbital angular momentum quantum numbers (smaller /) possess increasing numbers of nodes in their radial functions and are referred to as increasingly 'penetrating'. Thus, a 3s electron experiences a larger effective nuclear charge and is more tightly bound than a 3p electron; a 3p is in turn more tightly bound than a 3d. Next, we recall that the energy separations between adjacent principal quantum shells in hydrogen decrease with increasing n. Taking both factors together, we expect that sooner or later, with respect to increasing atomic number, the more tightly bound orbital subsets of the nih principal quantum shell will be more tightly bound and decrease in energy below the higher orbital subsets of the (n-l)th principal shell. 1.2 Complexes and Coordination Compounds 3 For neutral atoms, that cross-over begins around the start of the transition-metal series. The balance between the 4s and 3d orbital energies is delicate, however, and other factors, not discussed so far, can reverse the general trend. One such factor is the exchange stabilization associated with the filled and half-filled d shell. This will be familiar from discussion of ionization energies throughout the first long row of the periodic table when one considers the marked discontinuities at the p3 and p6 configurations; this theme is taken up in more detail in Chapter 8. Now consider the ionization process yielding the M2+ ions in the first row transition-metal series. The configuration adopted in the ion does not depend solely upon the relative orbital energies of the (energetically close) 4s and 3d orbitals in the neutral atom. It also depends upon the relative energies of the putative ions 3dn~24s2 and 3i/"4s°, for example. Let us consider each in turn. Removal of electrons from the 3d shell relieves some electron - electron repulsion and deshields the 4s orbital somewhat: both 3d and 4s shells will be more tightly bound in an M2+ ion. Removal of electrons from the 4s shell, however, depletes the inner (sub-nodal) regions of their electron density with the result that the 3d orbitals are very much less well shielded and become much more tightly bound. It is perfectly possible in principle, and actually the case in practice, that the 3d orbital energy dips down below that of the 4s orbital as a result.

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Transition Metal

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