d-electron elements have distinctive physical properties
that make them useful as structural metals and catalysts, and distinctive
electronic properties that find applications in the molecules of
living organisms. The physical properties will be discussed in this
section, and the electronic properties in the next.
The four series of transition metals, and the electron-filling process
for the first of these, are shown below. Generally, each new transition
metal along one row has two electrons in the outer s orbital,
and an increasing number of electrons in the buried d orbital
belonging to the preceding principal quantum number. In Row 4 chromium
(Cr) and copper (Cu) are exceptions to this rule, stealing one of
the two s electrons to half-fill or fill the d orbitals.
These are minor exceptions, and it is the filling principle that
In chemical reactions of the transition metals, the s electrons
are lost most easily, and +1 and +2 oxidation states are common.
Higher oxidation states also are possible for atoms that have d
electrons and, in principle, the highest possible state would correspond
to the loss of all of the outer s and d electrons:
+3 for scandium, +4 for titanium, and +7 for manganese. Electron
pairing complicates matters from iron onward, as we shall see.
Electron-shell diagram for transition metals.
The s orbital is the outer shell, and the d orbitals from the preceeding
principle quantum number are buried more deeply.
The A-group representative elements in the periodic table are interrupted
to accommodate the transition metals (B group).