Transition metals are famous for producing brilliantly colored ions in solution—deep blues, greens, purples, and reds. This behavior is central to IB Chemistry Topic 13 (HL) and is one of the key reasons transition metals are so important in biological systems, industry, and analytical chemistry. The color arises from the unique electronic structure of partially filled d-orbitals and the interaction between the metal ion and surrounding ligands.
What Makes Transition Metal Ions Colored?
Transition metal ions are colored because electrons in their partially filled d-orbitals absorb specific wavelengths of visible light to transition between split energy levels.
This phenomenon occurs only when:
- The metal has a partially filled d-subshell
- The metal ion forms a complex with ligands
- The d-orbitals are split into two different energy levels
When light hits the ion:
- Some wavelengths are absorbed
- Electrons jump from lower to higher energy d-orbitals
- The remaining transmitted or reflected wavelengths produce the visible color
The Role of d-Orbital Splitting
When ligands surround a transition metal ion, they create an electric field.
This field causes the five d-orbitals to split into two energy groups:
- A lower-energy set
- A higher-energy set
This splitting is called crystal field splitting (ΔE).
Key idea:
The size of the energy gap determines the wavelength of light absorbed.
For example:
- Larger energy gap → absorbs higher-energy (blue/violet) light → appears yellow/orange
- Smaller energy gap → absorbs lower-energy (red) light → appears blue/green
Why Partially Filled d-Subshells Are Required
Transition metals must have partially filled d-orbitals for colors to appear.
If the d-orbitals are completely empty (d⁰) or fully filled (d¹⁰), there are no electrons available for d–d transitions.
Examples:
- Sc³⁺ is colorless (d⁰)
- Zn²⁺ is colorless (d¹⁰)
- Cu²⁺ is colored (d⁹)
- Fe²⁺ and Fe³⁺ are colored (d⁶ and d⁵)
This is why zinc, cadmium, and mercury are not considered true transition metals under the IB definition—they form colorless ions.
How Ligands Affect Color
Different ligands cause different amounts of d-orbital splitting.
This changes the wavelength of light absorbed, leading to different colors.
Ligands that create large splitting:
- CN⁻
- NH₃
These often produce deeply colored complexes.
Ligands that create small splitting:
- H₂O
- F⁻
- Cl⁻
These tend to produce complexes with lighter or different hues.
Example:
[Cu(H₂O)₆]²⁺ → blue
[Cu(NH₃)₄(H₂O)₂]²⁺ → deep royal blue
The color change comes from ligand substitution altering the energy gap.
How Oxidation State Affects Color
Higher oxidation states usually cause:
- Greater attraction between metal and ligands
- Stronger splitting
- Different colors
Examples:
- Mn²⁺: pale pink
- MnO₄⁻ (Mn⁷⁺): deep purple
A dramatic shift in oxidation state creates large differences in ΔE, and therefore large differences in color.
How Coordination Number Influences Color
Coordination number describes how many ligands surround the metal ion.
Common numbers:
- 4 (tetrahedral or square planar)
- 6 (octahedral)
Different geometries alter crystal field splitting patterns:
- Octahedral complexes usually have a larger splitting
- Tetrahedral complexes have smaller splitting
This shifts the absorbed wavelength and the observed color.
Examples of Transition Metal Colors (IB Relevant)
- Cu²⁺ → blue
- Ni²⁺ → green
- Cr₂O₇²⁻ → orange
- CrO₄²⁻ → yellow
- Fe²⁺ → pale green
- Fe³⁺ → yellow-brown
- MnO₄⁻ → purple
- V³⁺ → green
- V²⁺ → violet
These colors often appear in lab work and exam questions.
Common IB Misunderstandings
“All metal ions are colored.”
False—only transition metal ions with partially filled d-orbitals produce colors.
“Color comes from electrons jumping between shells.”
Color comes from d–d transitions within the d-subshell, not shell-to-shell movement.
“The metal controls the color entirely.”
Ligands, geometry, and oxidation state also play major roles.
“All complexes of the same metal are the same color.”
Different ligands produce different colors.
FAQs
Why does the same metal ion appear different colors with different ligands?
Because ligand strength changes the d-orbital splitting and alters which wavelength is absorbed.
Why is a Zn²⁺ solution colorless?
Its d-orbitals are completely filled (d¹⁰), so no d–d transitions occur.
Why do transition metals form so many colorful compounds?
They have partially filled d-orbitals and easily form complexes with many ligands.
Conclusion
Transition metal ions are colored because their d-orbitals split into different energy levels when ligands surround them. Electrons absorb specific wavelengths of light to jump between these levels, and the remaining light gives the ion its color. Factors such as ligand type, oxidation state, and coordination number all influence the color. This makes transition metal chemistry one of the most visually distinctive and conceptually rich areas of the IB syllabus.
