Periodic Table - Transition Metal Chemistry - Doc Brown's Chemistry  Revising Advanced Level Inorganic Chemistry Periodic Table Revision Notes

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

Why are most transition metal complexes coloured? Why are some complexes colourless? What is the origin of colour? The electronic theory of 3d orbital splitting by the ligands to create the possibility of quantum level change is described and why many transition metal complexes absorb visible light so that the colour you see is the visible light that is transmitted i.e. unabsorbed. Note how changing the oxidation state or the ligand changes the complex electronically sufficiently to produce different colours even for the same central transition metal ion.

(c) doc b GCSE/IGCSE Periodic Table Revision Notes * (c) doc b GCSE/IGCSE Transition Metals Revision Notes

INORGANIC Part 10 3d block TRANSITION METALS sub–index: 10.1–10.2 Introduction 3d–block Transition Metals * 10.3 Scandium * 10.4 Titanium * 10.5 Vanadium * 10.6 Chromium * 10.7 Manganese * 10.8 Iron * 10.9  Cobalt * 10.10 Nickel * 10.11 Copper * 10.12 Zinc * 10.13 Other Transition Metals e.g. Ag and Pt * Appendix 1. Hydrated salts, acidity of hexa–aqua ions * Appendix 2. Complexes & ligands * Appendix 3. Complexes and isomerism * Appendix 4. Electron configuration & colour theory * Appendix 5. Redox equations, feasibility, Eø * Appendix 6. Catalysis * Appendix 7. Redox equations * Appendix 8. Stability Constants and entropy changes * Appendix 9. Colorimetric analysis and complex ion formula * Appendix 10 3d block – extended data * Appendix 11 Some 3d–block compounds, complexes, oxidation states & electrode potentials * Appendix 12 Hydroxide complex precipitate 'pictures', formulae and equations

Advanced Level Inorganic Chemistry Periodic Table Index * Part 1 Periodic Table history * Part 2 Electron configurations, spectroscopy, hydrogen spectrum, ionisation energies * Part 3 Period 1 survey H to He * Part 4 Period 2 survey Li to Ne * Part 5 Period 3 survey Na to Ar * Part 6 Period 4 survey K to Kr and important trends down a group * Part 7 s–block Groups 1/2 Alkali Metals/Alkaline Earth Metals * Part 8  p–block Groups 3/13 to 0/18 * Part 9 Group 7/17 The Halogens * Part 10 3d block elements & Transition Metal Series * Part 11 Group & Series data & periodicity plots * All 11 Parts have their own sub–indexes near the top of the pages

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

  • Transition metal ions can be identified by their colour.

    • The colour arises when some of the wavelengths of visible light are absorbed and the remaining wavelengths of light are transmitted or reflected, so you experience the net effect.

    • In transition metal species d electrons move from the ground state to an excited state when light is absorbed.

    • The energy difference between the ground state and the excited state of the d electrons is given by the equation:

      • ∆E = hν = hc/λ (more details on Planck's equation further down the page)

    • Changes in oxidation state, co-ordination number and ligand alter ∆E and this leads to a change in colour e.g. of the transition metal complex ion.

    • The absorption of visible light is used in spectroscopy e.g. using a simple colorimeter to determine the concentration of coloured ions in solution and also the chemical formula of a complex ion.

  • 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), hPlanck'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.

  • 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

  • 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.

Scandium * Titanium * Vanadium * Chromium * Manganese * Iron * Cobalt * Nickel * Copper * Zinc * Silver & Platinum

Introduction 3d–block Transition Metals * Appendix 1. Hydrated salts, acidity of hexa–aqua ions * Appendix 2. Complexes & ligands * Appendix 3. Complexes and isomerism * Appendix 4. Electron configuration & colour theory * Appendix 5. Redox equations, feasibility, Eø * Appendix 6. Catalysis * Appendix 7. Redox equations * Appendix 8. Stability Constants and entropy changes * Appendix 9. Colorimetric analysis and complex ion formula * Appendix 10 3d block – extended data * Appendix 11 Some 3d–block compounds, complexes, oxidation states & electrode potentials * Appendix 12 Hydroxide complex precipitate 'pictures', formulae and equations

 

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