We have seen the trade-off between orbital overlap and magnetism before (in Chapter 5) in the context of paramagnetic transition metal complexes. Depending on the way the spins order, metals and alloys in this part of the periodic table can be ferromagnetic (spins on neighboring atoms aligned parallel, as in the case of Fe or Ni) or antiferromagnetic (spins on neighboring atoms antiparallel, as in the case of Mn). At this point the elements become magnetic. For elements beyond V, the orbital overlap is so poor that the 3d electrons are no longer effective in bonding, and the valence electrons begin to unpair. We can explain this effect by remembering that the 3d orbitals are progressively contracting as more protons are added to the nucleus. The cohesive energy actually decreases going from V to Mn, even though the number of valence electrons is increasing.
#4d orbital series#
The 3d series has a "crater" in the cohesive energy plot where there was a peak in the 5d series. In the 3d series, we see the expected increase in cohesive energy going from Ca (4s 2) to Sc (4s 23d 1) to Ti (4s 23d 2) to V (4s 23d 3), but then something very odd happens. The 3d elements (Sc through Zn) are distinctly different from the 4d and 5d elements in their bonding (and consequently in their magnetic properties). The heat of vaporization (the cohesive energy) of metals in the 3d and 5d series, measured at the melting point of the metal. We do not see magnetism in the 4d or 5d metals or their alloys because orbital overlap is strong and the bonding energy exceeds the electron pairing energy. Because of their strong bonding energy, elements in the middle of the 4d and 5d series have very high melting points. For example, Pt metal must be promoted from the 6s 15d 9 atomic ground state to 6s 15d 76p 2 in order to make six bonds per atom, and the energy cost of promoting electrons from the 5d to the 6p orbitals is reflected in the net bonding energy. Elements past Mo and W have more d electrons, but some of them are spin paired and so some promotion energy is needed to prepare these electrons for bonding. Mo and W have the most bonding energy because they can use all six of their valence electrons in bonding without promotion. The number of bonding electrons, and therefore the bonding energy, increases steadily going from Rb to Mo in the 4d series, and from Cs to W in the 5d series. number of valence electrons (below left) has a "volcano" shape that is peaked at the elements Mo and W (5s 14d 5 and 6s 15d 5, respectively).
In the 4d and 5d series, a plot of cohesive energy vs. The 5d orbitals are well shielded by the complete n=4 shell so they have good orbital overlap with neighboring atoms. The increased nuclear charge is felt most strongly by the 6s, which takes on the character of an inert electron pair. In the third transition series, the situation is reversed. Therefore the 4s and 4p orbitals are more effective in bonding than the 3d.
In the first transition series, shielding of the 3d orbitals is poor. Schematic representation of the sizes of different orbitals of Cr and W. Weak overlap of 3d orbitals gives narrow d-bands and results in the emergence of magnetic properties as discussed below. Past V in the first row of the transition metals, the 3d electrons become much less effective in bonding because they overlap weakly with their neighbors. In the 3d series, the contraction of orbitals affects the ability of the d electrons to contribute to bonding. These trends explain the distinct behavior of the 3d elements relative to those in the 4d and 5d series. Moving down the periodic table ( V - Nb - Ta ), the d orbitals expand because of the increase in principal quantum number. As we move across the periodic table ( Sc - Ti - V - Cr - Fe ), the d orbitals contract because of increasing nuclear charge. There are two important periodic trends that are related to orbital size and orbital overlap. While electrons in s and p orbitals tend to form strong bonds, d-electron bonds can be strong or weak.
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