Formation of Coloured Ions (AQA A Level Chemistry): Revision Note

Colour in Transition Metal Ions

• Most transition metal complexes are coloured

• The colour arises because transition metal ions absorb specific wavelengths of visible light due to d–d electronic transitions

• When white light passes through a coloured complex solution, certain wavelengths are absorbed

  • The colour observed is the complementary colour of the light absorbed — that is, the colour corresponding to the wavelengths that are transmitted or reflected

• For example, aqueous copper(II) ions absorb light from the red region of the visible spectrum

  • The complementary colour observed is blue-green (cyan)

Electron promotion in transition metal ions

• Electron promotion explains the colour of transition metal complexes

• In an isolated transition metal ion (with no ligands attached), the five 3d orbitals are equal in energy

  • Orbitals with the same energy are described as degenerate

• When ligands bond to the central metal ion by forming dative covalent bonds, the 3d orbitals split into two sets with different energies

  • These are described as non-degenerate orbitals

  • The energy difference between the two sets is called ΔE

• When light passes through a solution containing a transition metal complex, an electron can absorb energy equal to ΔE and become promoted from a lower-energy 3d orbital to a higher-energy 3d orbital

  • This process is known as electron promotion (or excitation)

• The energy absorbed is given by:

h = Planck's constant (6.626 x 10^-34 m² kg s^-1)

= frequency (Hertz, Hz or s^-1)

• The remaining wavelengths of light that are not absorbed combine to produce the complementary colour that is observed

• For example, in an octahedral nickel(II) complex, the 3d orbitals split into two energy levels, and electron promotion between these levels results in the observed colour

Changes in Colour

• Transition metal complexes absorb light with a frequency that corresponds exactly to the energy gap (ΔE) between the split, non-degenerate d orbitals

• The wavelengths of light that are not absorbed are transmitted

  • These combine to produce the complementary colour that is observed

• The size of the energy gap (ΔE) is influenced by several factors, including:

  • The type of ligand (different ligands cause different amounts of splitting)

  • The coordination number (which affects the geometry of the complex)

  • The oxidation state of the metal ion (higher oxidation states generally increase ΔE)

• Changes in these factors alter ΔE, and therefore change the wavelength of light absorbed and the colour observed

Type of ligand

• Different ligands cause different amounts of splitting of the d orbitals in a transition metal complex

• The extent of splitting depends on the strength of interaction between the ligands and the metal ion (often described in terms of ligand field strength)

  • Stronger-field ligands produce a larger energy gap (ΔE)

• If ΔE changes, the frequency of light absorbed also changes

  • As a result, a different wavelength of light is absorbed, and a different complementary colour is observed

• This means that complexes containing the same transition metal ion in the same oxidation state can have different colours if the ligands are different

• For example, the hexaaquacopper(II) ion, [Cu(H₂O)₆]²⁺, is light blue

  • When ammonia ligands replace some of the water ligands to form [Cu(NH₃)₄ (H₂O)₂]²⁺, the solution becomes dark blue

• In both complexes, copper has an oxidation state of +2

  • The difference in colour shows that the ligands surrounding the metal ion affect the size of ΔE and therefore the colour of the complex

Coordination number

• The coordination number affects the strength of interaction between the metal ion and its ligands

  • Different coordination numbers often result in different geometries (for example, octahedral or tetrahedral), which alter the splitting of the d orbitals and therefore the size of ΔE

• In practice, changing the coordination number usually also involves changing the ligand, so both factors influence the strength of the metal–ligand interaction

Oxidation State

• The strength of attraction between a transition metal ion and the lone pairs on the ligands depends on the effective nuclear charge of the metal ion

• For example, manganese(II) and iron(III) both have the electron configuration [Ar]3d⁵

  • However, Fe³⁺ has a higher nuclear charge than Mn²⁺

  • This results in a stronger attraction between the metal ion and the ligands in [Fe(H₂O)₆]³⁺ compared with [Mn(H₂O)₆]²⁺

• As a result, the splitting of the d orbitals (ΔE) is larger for [Fe(H₂O)₆]³⁺

  • The larger ΔE means that higher-energy (shorter wavelength) light is absorbed

• For example:

  • [Mn(H₂O)₆]²⁺ appears pale pink because it absorbs light in the green region of the spectrum

  • [Fe(H₂O)₆]³⁺ absorbs higher-energy blue light and therefore appears orange (the complementary colour)

• Similarly, when comparing iron(II) and iron(III):

  • [Fe(H₂O)₆]²⁺ absorbs light in the red region and appears pale green

  • [Fe(H₂O)₆]³⁺ absorbs in the blue region and appears orange

• This shows that a higher oxidation state generally leads to stronger metal–ligand interactions, a larger ΔE, and a change in the colour of the complex

Visible Light Spectroscopy

• Spectroscopy is the study of the interaction between electromagnetic radiation and matter

• A simple colorimeter is used to determine the concentration of coloured solutions, such as transition metal ion solutions

• The colorimeter passes light of a selected wavelength through the solution

  • The wavelength is chosen using coloured filters, and the intensity of transmitted light is measured by a detector

• The filter selected corresponds to the wavelength of light that is most strongly absorbed by the solution

  • This is the complementary colour to the observed colour of the solution

• For example, a blue solution absorbs light from the red region of the spectrum

  • Therefore, a red filter is used so that red light passes through the solution, and maximum absorption occurs

• Some of the light is absorbed by the solution, and the remaining light passes through to the detector

  • The amount of light absorbed can then be used to calculate the concentration of the solution

• To determine the concentration of a coloured solution, a calibration curve must first be constructed

• This involves measuring the absorbance of a series of standard solutions with known concentrations

• At low concentrations, absorbance is directly proportional to concentration

  • This relationship is known as the Beer–Lambert law, and it produces a straight-line calibration graph

• Once the calibration curve has been plotted, the absorbance of an unknown solution of the same complex ion can be measured.

  • The concentration is then determined by locating the absorbance value on the y-axis, drawing across to the calibration line, and then down to the concentration axis

• Limitations of visible spectroscopy include:

  • Very dark solutions may absorb too much light, making accurate measurement difficult

  • Very pale solutions may have absorbance values close to the sensitivity limit of the colorimeter

• One way to overcome low absorbance is to convert the complex into a more intensely coloured species using ligand exchange

• For example, adding thiocyanate ions (SCN⁻) to aqueous iron(III) ions forms a blood-red complex:

                    [Fe(H₂O)₆]³⁺   (aq)  + SCN^- (aq) ⇌  [Fe(H₂O)₅ SCN]²⁺ (aq) + H₂O (l)

    pale orange           colourless            blood-red complex