Appendix 4. Electron configurations and the theory and variation of complex ion colour

• For the 3d-block know the complete order of filling of the sub-shells from Z=21 to 30 and be able to write out the full or abbreviated electron configuration.

  • See under each element and even more detail in Electron Configurations in Periodic Table section 2.2

  • Transition metal electron arrangements are listed

  • The number of orbitals per sub-shell, 1 for s, 3 for p, and 5 for d sub-shell.

  • PLEASE watch out for the two ‘quirks’ for Cr and Cu atoms and the order of electron removal when forming positive ions e.g. for the 3d block of transition metals, you remove the 4s electrons first, before any of the 3d electrons.

  • The idea of atomic orbitals as the space/shape of a particular electronic level or sub-shell helps ion this section.

• Transition metals can be identified by the colour of their complexes which of course is a very characteristic feature of their chemistry (e.g. the hydroxide precipitates which are, of course, all neutral complexes).

• The colour can varies with change in (i) oxidation state, (ii) ligand and (iii) co-ordination number or shape (which in turn depends on the ligand and oxidation state) and obviously changing the transition metal itself, will give another range of differently coloured compounds.

• All of these factors are linked to the electronic state of the central metal ion, so, if the electronic levels are changed by change in oxidation state or ligand, the difference between quantum levels changes, therefore the wavelength of the light photons absorbed changes, i.e. the observed colour changes e.g.

  • e.g. (i) The same ligand (H₂O), shape and co-ordination number but different oxidation state.

    • and

    • [Fe(H₂O)₆]²⁺, pale green iron(II) ion and [Fe(H₂O)₆]³⁺, yellow-brown iron(III) ion.

      • Oxidation states +2 and +3, both octahedral complexes with co-ordination number 6.

  • e.g. (ii) The same oxidation state, shape and co-ordination number but different ligand.

    • and

    • [Ni(H₂O)₆]²⁺, green hexaaquanickel(II) ion and [Ni(NH₃)₆]²⁺, pale blue hexaamminenickel(II) ion.

      • Both oxidation state +2, both octahedral complexes with co-ordination number 6, but different ligands i.e. water and ammonia.

  • e.g. (iii) The same oxidation state but with a different ligand, shape and co-ordination number.

    • and

    • [Cu(H₂O)₆]²⁺, blue hexaaquacopper(II) ion and [CuCl₄]^2-, yellow tetrachlorocuprate(II) ion.

    • Same oxidation state +2, but different ligands (water and chloride ion), different shape (octahedral and tetrahedral) and different co-ordination number (6 and 4).

• COLOUR THEORY for transition element complexes: The argument is presented from the point of view of an octahedral complex, but similar arguments apply for a tetrahedral or square planar complex.

• There are five 3d sub-shell orbitals whose 3D spatial representations are shown below. Theoretically it is considered that the ligands in an octahedral complex approach the central metal ion along the x, y and z axis, which would minimise the repulsion between the orbitals of bonding electrons in the six M-ligand dative covalent bonds (note that 4s and 4p orbitals are involved in complex ion bonding).

  • The d orbitals point either along or between the x,y,z Cartesian axes.

• The approach and bonding of these ligands raises the energy of all of the 3d orbitals, but not all equally so.

  • For an octahedral complex, the two orbitals lying on the x, y and z axes (4) and (5) experience more repulsion than the other three orbitals lying between the x, y and z axes (1), (2) and (3) when the six co-ordinate covalent ligand - metal ion bonds are formed.

  • This unequal ligand repulsion causes a splitting in the 3d orbital quantum levels.

  • In each of the four box diagrams (1)-(4) below, the five raised 3d 'degenerate' (meaning equal) orbitals are shown on the left, and the 'splitting' effect of the ligands is shown on the right.

  • The lower three 3d orbitals represent the 'new' ground state.

  • The upper two 3d orbitals represent either an upper ground state if the lower 3d levels are fully occupied, or more pertinent to colour theory, a potentially upper excited state if they are not fully occupied.

  • We can now consider what possible electronic 'transitions' can take place for four different ions - coloured and colourless.

The electronic ground states of scandium(III), titanium(III), copper(II) and zinc(II) are illustrated below.

The electronically excited states of titanium(III) and copper(II) are illustrated below.

• The colour arises from electronic transitions from the ground state to excited states, the energy needed can be calculated using

  • Planck's Equation, ΔE = hv , E = energy of a single photon (J), h =  Planck's Constant (6.63 x 10-34 JHz^-1), v = frequency (Hz).

  • Therefore if the photo energy/frequency is equal to ΔE  then energy is absorbed and an electron can be promoted from the lower 3d level to the higher 3d level.

  • If ΔE is in the visible light frequency range the complex will be 'coloured'.

  • In the case of coloured transition metal complexes, the colour arises from visible light energy absorption on promoting electrons from the lower 3d levels to the higher 3d levels.

• However, this can only occur if there is at least one electron in the 'lower' 3d orbitals and at least one half-filled 'higher' 3d quantum level, i.e. the minimum pre-conditions for an electronic transition or 'excitation'.

• Consequently because there a lack of such possible transitions in Sc(III) and Zn(II) their compounds are usually colourless i.e. no light absorbed in the visible region of the spectrum

• In the true transition metals from Ti to Cu, it is possible for the electromagnetic radiation energy to produce this excitation from the lower to the higher 3d sub-levels and it is usually in the visible region.

• Certain frequencies-frequency ranges of visible radiation are absorbed and the perceived colour arises from the frequencies not absorbed i.e. the transmitted visible light.

• The electronic structure and colour of some typical 'simple' aqueous ions is shown below. They are all hexa-aqua ions of an octahedral shape except ...

  • copper(I) cannot form a stable simple Cu⁺_(aq) ion, but copper(I) compounds tend to be colourless when pure e.g CuCl, copper(I) chloride,

  • but copper(II) forms the blue square planar [Cu(H₂O)₄]²⁺ ion.

• The colour you see in a transition metal compound is the visible light that isn't absorbed by the 3d electronic transitions. For example, copper(II) complexes often absorb in the red area of the visible spectrum, so the resulting colour observed is green-blue.

• Colour changes in transition metal reactions can arise from change of ligand, change in co-ordination number or change in metal oxidation state (sometimes several of these simultaneously.

  • See cobalt(II) ion reactions with ammonia + oxygen or chloride ion).

• The colours are quite useful for simple transition metal ion identification tests e.g. precipitates with sodium hydroxide and ammonia (see pictures) and the thiocyanate test for iron(III) ions.

• Ultraviolet and visible absorption spectra

  • x-ref complimentary colour table in Appendix 9. Colorimetry - be able to suggest a colour from the absorption spectrum.

  • need visible spectra pictures

• Dyes and pigments

  • Monastral blue

  • other porphyrin complexes for dyes in paint

• Ultraviolet and visible spectroscopy can be used to determine the concentration of metal ions in solution, usually after the addition of a suitable ligand to intensify the colour using the more elaborate technique of spectrophotometry or the simpler technique of colorimetry - appendix 9.

  • Theory: diagram, spectra

  • Examples:

• Colorimetric analysis of coloured solutions for quantitative analysis using a colorimeter is described in Appendix 9.