Transition Elements, Complex Ions and Colour

(HL) What counts as a transition element

A transition element forms at least one stable ion with a partially filled d sub-shell. That definition excludes $\ce{Sc}$ (its only ion $\ce{Sc^{3+}}$ is $\text{3d}^0$) and $\ce{Zn}$ ($\ce{Zn^{2+}}$ is $\text{3d}^{10}$), even though both sit in the d block.

Because the 3d and 4s sub-shells are very close in energy, electrons from both can be involved in bonding — hence the variable oxidation states. Manganese runs from $+2$ to $+7$.

Complex ions

A ligand is a species with a lone pair that forms a coordinate bond to the metal ion. Common ligands are $\ce{H2O}$, $\ce{NH3}$, $\ce{Cl-}$ and $\ce{CN-}$; the usual coordination number is $6$ (octahedral), sometimes $4$.

Why the complexes are coloured

In a free ion the five d orbitals have the same energy. A set of ligands splits them into two groups separated by an energy gap $\Delta$. If $\Delta$ falls in the visible range, a d electron can absorb a photon and jump the gap — a d–d transition. The colour you see is the complementary colour of the light absorbed: a complex that absorbs red looks green.

So a complex is coloured only if the d sub-shell is partially filled: $\text{3d}^0$ has no electron to promote and $\text{3d}^{10}$ has no vacancy to promote into. That is exactly why $\ce{Sc^{3+}}$ and $\ce{Zn^{2+}}$ solutions are colourless.

$\Delta$ — and therefore the colour — changes with the ligand (spectrochemical series $\ce{CN- > NH3 > H2O > Cl-}$), with the oxidation state of the metal, and with the geometry.