Transition metals form colorful ions, complexes, and compounds. The colors are characteristic of the element and whether it’s in aqueous solution or another solvent besides water. The colors are helpful in qualitative analysis because they offer a clue to sample composition. Here is a look at transition metal colors in aqueous solution and an explanation of why they occur.
Why Transition Metals Form Colored Complexes
Transition metals form colored solutions and compounds because these elements have unfilled d orbitals. The metal ions aren’t actually colored on their own because the d orbitals are degenerate. In other words, they all have the same energy, which corresponds to the same spectral signal. When transition metal ions form complexes and compounds with other molecules, they become colored. A complex forms when a transition metal bonds to one or more neutral or negatively charged nonmetals (ligands). The ligand changes the shape of d orbitals. Some of the d orbitals gain a higher energy than before, while others move to a lower energy state. This creates an energy gap. The wavelength of the photon that is absorbed depends on the size of the energy gap. (This is why splitting of s and p orbitals, while it occurs, does not produce colored complexes. Those gaps would absorb ultraviolet light and not affect the color in the visible spectrum.)
Unabsorbed wavelengths of light pass through a complex. Some light is also reflected back from a molecule. The combination of absorption, reflection, and transmission results in the apparent colors of the complexes. For example, an electron may absorb red light and become excited into a higher energy level. Since the non-absorbed light is the color reflected, we would see a green or blue color.
Complexes of a single metal may be different colors depending on the oxidation state of the element.
Why Not All Transition Metals Display Colors
But, not all oxidation states produce colors. A transition metal ion with zero or ten d electrons forms a colorless solution.
Another reason not all elements in the group display colors is that they aren’t all technically transition metals. If an element must have an incompletely filled d orbital to be a transition metal, then not all d block elements are transition metals. So, zinc and scandium aren’t transition metals under the strict definition because Zn2+ has a full d level, while Sc3+ has no d electrons.
Transition Metal Ion Colors in Aqueous Solution
Here is a table of common transition metal ion colors in aqueous solution. Use this as an aid for AP Chemistry and qualitative analysis, especially in conjunction with other diagnostic tools, such as the flame test.
| Transition Metal Ion | Color |
| Ti2+ | Pale Brown |
| Ti3+ | Purple |
| V2+ | Purple |
| V3+ | Green |
| V4+ | Blue-Gray |
| V5+ | Yellow |
| Cr2+ | Blue-Violet |
| Cr3+ | Green |
| Cr6+ | Orange-Yellow |
| Mn2+ | Pale Pink |
| Mn7+ | Magenta |
| Fe2+ | Olive Green |
| Fe3+ | Yellow |
| Co2+ | Red to Pink |
| Ni2+ | Bright Green |
| Cu2+ | Blue-Green |
Other Transition Metal Complex Colors
The colors of transition metal complexes often vary in different solvents. The color of the complex depends on the ligand. For example, Fe2+ is pale green in water, but forms a dark green precipitate in a concentrated hydroxide base solution, carbonate solution, or ammonia. Co2+ forms a pink solution in water, but a blue-green precipitate in hydroxide base solution, straw-colored solution in ammonia, and pink precipitate in carbonate solution.
Elements belonging to the lanthanide series also form colored complexes. The lanthanides are also known as the inner transition metals or simply as a subclass of the transition metals. However, the colored complexes are due to 4f electron transitions. The colors of the lanthanide complexes are not as influenced by the nature of their ligand and are pale compared to transition metal complexes.
References
- Cotton, F. Albert; Wilkinson, Geoffrey; Murillo, Carlos A.; Bochmann, Manfred (1999). Advanced Inorganic Chemistry (6th ed.). New York: Wiley-Interscience. ISBN 0-471-19957-5.
- Harris, D.; Bertolucci, M. (1989). Symmetry and Spectroscopy. Dover Publications.
- Huheey, James E. (1983). Inorganic Chemistry (3rd ed.). Harper & Row. ISBN 0-06-042987-9.
- Levine, Ira N. (1991). Quantum Chemistry (4th ed.). Prentice Hall. ISBN 0-205-12770-3.