The unique structure of the transition metals causes them to form brightly colored compounds. This structure affects the way light is absorbed, transmitted and reflected. The oxidation state of the particular element affects the colors of the compounds it forms.
Electrons at the d orbital affect the color of transition metal compounds. Therefore, different electron bonds in molecules allow manganese, for example, to form compounds ranging from dark purple to pale pink. These 5d electrons become more filled as one moves from the left to the right on the periodic table. Since the d orbitals are filled in zinc, it forms nearly colorless compounds.
Electrons absorb light of a certain wavelength to ascend to the next orbital, and the human eye sees the wavelengths that are not absorbed. Therefore, the energy gap between the higher and lower orbital levels is ultimately responsible for the variation in colors.
Transition metals have many common properties in addition to forming these highly colored compounds. They are all low ionization energy and have positive oxidation states. Transition metals have a tendency to be very hard yet remain malleable. They have high melting and boiling points. In addition, their high electrical conductivity makes transition metals ideal for use in electrical semiconductors.