6.2.3 Coloured Ions

Perception of colour

• Most transition metal compounds appear coloured. This is because they absorb energy corresponding to certain parts of the visible electromagnetic spectrum
• The colour that is seen is made up of the parts of the visible spectrum that aren’t absorbed
• For example, a green compound will absorb all frequencies of the spectrum apart from green light, which is transmitted
• The colours absorbed are complementary to the colour observed

The colour wheel showing complementary colours in the visible light region of the electromagnetic spectrum

• Complementary colours are any two colours which are directly opposite each other in the colour wheel
  • For example, the complementary colour of red is green and the complementary colours of red-violet are yellow-green

Splitting of 3d energy levels

• In a transition metal atom, the five orbitals that make up the d-subshell all have the same energy.
• Ions that have completely filled 3d energy levels (such as Zn²⁺) and ions that have no electrons in their 3d subshells (such as Sc³⁺) are not coloured
• Transition metals have a partially filled 3d energy level
• When ligands attach to the central metal ion the energy level splits into two levels with slightly different energies
  • If one of the electrons in the lower energy level absorbs energy from the visible spectrum it can move to the higher energy level
  • This process is known as promotion / excitation
• The amount of energy absorbed depends on the difference between the energy levels
  • A larger energy difference means the electron absorbs more energy
• The amount of energy gained by the electron is directly proportional to the frequency of the absorbed light and inversely proportional to the wavelength

Upon bonding to ligands, the d orbitals of the transition element ion split into sets of orbitals with different energies

The size of the splitting energy ΔE in the d-orbitals is influenced by the following four factors:

• The size and type of ligands
• The nuclear charge and identity of the metal ion
• The oxidation state of the metal
• The shape of the complex

The large variety of coloured compounds is a defining characteristic of transition metals

Size and type of ligand

• The nature of the ligand influences the strength of the interaction between ligand and central metal ion
• Ligands vary in their charge density
• The greater the charge density; the more strongly the ligand interacts with the metal ion causing greater splitting of the d-orbitals
• The further it is then shifted towards the region of the spectrum where it absorbs higher energy
• As a result, a different colour of light is absorbed by the complex solution and a different complementary colour is observed
• This means that complexes with the same transition elements ions, but different ligands, can have different colours
  • For example, the [Cu(H₂O)₆]²⁺ complex has a light blue colour
  • Whereas the [Cu(NH₃)₄(H₂O)₂]²⁺ has a dark blue colour despite the copper(II) ion having an oxidation state of +2 in both complexes

Ligand exchange of the water ligands by ammonia ligands causes a change in colour of the copper(II) complex solution

Oxidation number

• When the same metal has a higher oxidation number that will also create a stronger interaction with the ligands
• If you compare iron(II) and iron (III):
  • [Fe(H₂O)₆]²⁺ absorbs in the red region and appears green
  • But, [Fe(H₂O)₆]³⁺ absorbs in blue region and appears orange

Coordination number

• The change of colour in a complex is also partly due to the change in coordination number and geometry of the complex ion
• The splitting energy, ΔE, of the d-orbitals is affected by the relative orientation of the ligand as well as the d-orbitals
• Changing the coordination number generally involves changing the ligand as well, so it is a combination of these factors that alters the strength of the interactions