The dominant electronic valence configurations of atoms in chemical substances of a transition element of group "G" in period "n" is ("n" - 1)d[superscript "G"]"n"s[superscript 0]. Transition-metal chemistry is d orbital chemistry. In contrast, the ground states of free, unbound atoms derive, in most cases, from configurations ("n" - 1)d[superscript "G"-1]"n"s[superscript 1] or ("n" - 1)d[superscript "G"-2]"n"s[superscript 2]. Five features must be considered to resolve this paradox. Point (i), the d-s orbital-energy distance in relation to the averaged differences of dd, ds, and ss two-electron repulsion energies, has often been discussed in the literature. However, four equally important features are (ii) the different two-electron repulsion energies within the d shell; (iii) the spin-orbit coupling energies in the d shell;and, in particular, (iv) the "d-orbital collapse" below the s level for increasing nuclear charge around group 3 and (v) the different perturbations of the valence ("n" - 1)d and "n"s orbitals when a free atom enters an alloy or compound. There is a conceptual difference between spectroscopic ground "states" and chemically dominant ground "configurations". The ground states of unbound atoms mentioned in most chemistry textbooks have little meaning in chemistry. (Contains 2 tables, 3 figures and 3 notes.)