Formation of Coloured Ions (Oxford AQA International A Level Chemistry)

Colour in Transition Metal Ions

• Most transition element complexes are coloured

• Often, the transition metal ions can be identified by their colour

Colours of transition metal complexes

• A transition element complex solution which is coloured, absorbs wavelengths of light in the visible light region of the electromagnetic spectrum

• The observed colour is the complementary colour which is made up of wavelengths of light that are transmitted or reflected

  • For example, copper(II) ions absorb light from the red / orange end of the spectrum

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

Complementary colour wheel

Electron promotion

• An isolated transition element is one which is not bonded to any ligands

• In an isolated transition element ion, all of the 3d orbitals are degenerate

  • This means that they are equal in energy

• When ligands bond co-ordinately to the central metal ion, the 3d orbitals are split into two sets of non-degenerate orbitals

• The difference in energy between these two sets of non-degenerate orbitals is ΔE

• This amount of energy absorbed can be worked out from:

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

  • v = frequency (Hertz, Hz or s^-1)

• The equation for this relationship is:

ΔE = h x v

• If the frequency is not known, then the amount of energy absorbed can be worked out from:

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

  • c = velocity of light

  • = wavelength (cm^-1)

• The equation for this relationship is:

ΔE = h x c

• An electron starts in the ground state, which is a lower energy 3d non-degenerate orbital

• When light shines on a solution containing a transition element complex, an electron will absorb a specific amount of energy, ΔE

• The electron is now in the excited state, which is a higher energy 3d orbital

• This process is electron promotion

Electron promotion in an octahedral Ni²⁺ complex

• The amount of energy absorbed is linked to frequencies of light

• The frequencies of light which are not absorbed combine to make the complementary colour

Changes in Colour

• Transition element complexes absorb the frequency of light which corresponds to the exact energy difference (ΔE) between their non-degenerate d orbitals

• The frequencies of light which are not absorbed combine to make the complementary colour of the complex

• It is the complementary colour which is transmitted and observed

• However, the exact energy difference (ΔE) is affected by factors such as

  • the type of ligand

  • the coordination number

  • the oxidation state of the transition metal ion

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

• This causes greater splitting of the d-orbitals and means that more energy is required to promote an electron

• 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 element ions, but different ligands can have different colours:

  • [Cu(H₂O)₆]²⁺ complex has a light blue colour

  • [Cu(NH₃)₄ (H₂O)₂]²⁺ has a dark blue colour

  • However, the copper(II) ion has an oxidation state of +2 in both complexes

  • This is evidence that the ligands surrounding the complex ion affect the colour of the complex

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

Oxidation State

• The strength of the attraction between the transition metal ion and the electrons pairs in the dative covalent bonds can vary depending on the effective nuclear charge of the metal ion

• Manganese(II) and iron(III) have the same electronic configuration, 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵

  • Aqueous [Mn(H₂O)₆]²⁺

    • This absorbs energy in the green region of the spectrum

    • Therefore, it appears pink, following the principle of complementary colours

  • Aqueous [Fe(H₂O)₆]³⁺

    • It has a higher effective nuclear charge

    • This means it has a stronger interaction with the ligands

    • So, it absorbs in the blue (higher energy) region of the spectrum

    • Therefore, it appears orange

• When the same metal is in a higher oxidation state that will also create a stronger interaction with the ligands

  • Aqueous [Fe(H₂O)₆]²⁺

    • This absorbs energy in the red region of the spectrum

    • Therefore, it appears green

  • Aqueous [Fe(H₂O)₆]³⁺

    • It has a higher effective nuclear charge and stronger interactions with the ligands

    • So, it absorbs in the blue (higher energy) region of the spectrum and appears orange

Visible Light Spectroscopy

• Spectroscopy uses the absorption of visible light to find the concentration of coloured transition metal ion solutions

• The colorimeter passes different frequencies of light through a sample of the solution

  • The frequency absorbed is determined by the use of coloured filters and a detector

  • A filter is chosen matching the part of the spectrum where the absorption is strongest

  • This means choosing a filter that is the complementary colour to the colour of the solution

  • For example, a blue solution will absorb red, and therefore a red filter is used, so that only red light passes through the solution and maximum absorption occurs

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

Overview of visible light spectroscopy

• To determine the concentration of a coloured solution first a calibration curve has to be made

• This involves measuring the absorption of a set of standard solutions whose concentration is known

• The values are then plotted on a graph of absorption against concentration

Example calibration curve

• At lower concentrations the absorbance is directly proportional to the concentration of the coloured species so a straight line graph is obtained; this is known as the Beer-Lambert Law

• Once the calibration curve is plotted then the absorbance of any unknown solution of the same complex ion can be measured

• You can then deduce the concentration of the unknown sample by extrapolating from the absorbance to the calibration curve and down to the concentration