Coloured Ions (Edexcel A Level Chemistry): Revision Note

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

Changes in Colour

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 greater the splitting, the more energy is needed to promote an electron to the higher group of orbitals

• Therefore, the further it is shifted towards the region of the spectrum where it absorbs higher energy

  • As splitting increases, the light absorbed will shift away from the red end of the spectrum (longer wavelengths), towards the yellow end (shorter wavelengths)

• 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