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

 Appendix 11 3d–block compounds, complexes, oxidation states & electrode potentials

A summary table of some common complexes of the 3d–block transition metals (mainly the aqua complex ions, oxo–cations and oxy–anions) and an electrode potential chart of Sc to Zn

 

(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 11 A summary of some 3d–block compounds, complexes, oxidation states and electrode potentials

Most are mentioned in the detailed individual element notes, but some have been added to illustrate other oxidation states you may not encounter on your course – but some good oxidation number practice!

Ox. State Sc Ti V Cr Mn Fe Co Ni Cu Zn
+1, (I) CuI white(s)

[CuCl3]2–

+2, (II) [Ti(H2O)6]2+ violet(aq) [V(H2O)6]2+ violet(aq) MnO (s)

[Mn(H2O)6]2+ very pale pink(aq)

[Fe(H2O)6]2+ pale green(aq) CoO (s)

[Co(H2O)6]2+ pink(aq)

NiCl2 (s)

[Ni(H2O)6]2+ green(aq)

[Ni(CN)4]2–

[Cu(H2O)6]2+ blue green(aq) ZnO, ZnCO3 white(s)

[Zn(H2O)4]2+ colourless(aq)

+3, (III) Sc2O3 Sc(OH)3 white(s)

[Sc(H2O)6]3+ colourless(aq)

[Ti(H2O)6]3+ purple(aq) [V(H2O)6]3+ green(aq) Cr2O3 (s)

[Cr(H2O)6]3+ green(aq)

Mn2O3 brown(s) Fe2O3 brown(s)

[Fe(H2O)6]3+ yellowish–brown(aq)

[Co(NH3)6]3+(aq)
+4, (IV) TiO2 white(s)

[TiO]2+ colourless(aq)

TiCl4 colourless(l)

[VO]2+ blue(aq) MnO2 black(s)
+5, (V) V2O5 white(s)

VO43–

[VO2]+ yellow(aq)

+6, (VI) CrO3 (s)

Cr2O72– orange(aq)

CrO42– yellow(aq)

MnO42– green(aq) FeO42– (in s)
+7, (VII) KMnO4 dark purple(s)

MnO4 purple(aq)

  • Notes

    1. See REDOX pages for the meaning of oxidation state and how to work it out in a compound.

    2. Can you see in each case why the oxidation state is as quoted? i.e. can you work out the oxidation number of the 3d–block metal.

    3. The text is small to fit the table on a minimum of a 1024 x 768 screen.

    4. Nice pattern of maximum oxidation state from Sc to Mn i.e. equivalent to using/losing all the outer electrons (3dx 4sy) beyond the [Ar] core.

    5. All except scandium (Sc3+) form an M2+ ion.

    6. All except zinc form compounds with a (III) oxidation sate compound.

    7. Advanced Inorganic Chemistry Page Index and LinksOnly copper has important compounds of oxidation state +1.

 Standard Electrode Potential Chart Diagram for the 3d–block elements

  • Redox potential chart comments:

  • All data quoted is for standard conditions i.e. 298K, 1 atm. pressure and 1 mol dm–3 solutions of ions.

  • Other than the solid metals, MnO2 and FeO42–, hydrogen gas, you can assume all ions are in aqueous media.

  • Unless an oxyanion, oxocation or another ligand in a complex is indicated, you assume you are dealing with hexaaqua–metal ions (H2O ligand only).

  • Further comments below draw out some general patterns and other points of interest.

    • All except scandium (Sc3+), which is not that reactive to acids despite the relatively negative M/M3+ potential, form a hydrated M2+ ion either by reaction of the metal with acid or reduction of a higher oxidation state complex–compound.

    • The stable oxidation states in aqueous solution containing dissolved oxygen from air are for the hydrated ions ...

      • (only Sc3+), [TiO]2+, VO2+, Cr3+, Mn2+, Fe3+, Co2+ and Ni2+ (only Zn2+).

      • On the basis of the electrode potential chart above, the argument is simple. In neutral or acid solution the oxidising potential of the oxygen–proton–water system is +1.23V. Therefore any e.g. M3+/M2+ potential less positive than +1.23V will result in the oxidation of the lower oxidation state species to the higher oxidation state species in the presence of dissolved oxygen which is reduced to water.

        • Oxidation states higher than the stable ones tend to oxidise water liberating oxygen and as mentioned above, lower oxidation states tend to be reducing and liberate hydrogen from water.

        • So the Mn3+/Mn2+ and Co3+/Co2+ potentials lie above +1.23V so Mn3+ and Co3+ will oxidise water and cannot be stable in acid solution.

      • Note that the +4 oxidation states of Ti and V exist as hydrated oxo–cations because the high polarising power of the highly charged central metal ion causes deprotonation (see Appendix 1. Acidity of hexa–aqua ions).

      • The rest are [M(H2O)n]2+/3+ where n is usually 6, can be 4 for Cu and Zn.

    • Apart from iron, there is a tendency for the lower oxidation state to become increasingly more stable with increasing atomic number.

    • Higher oxidation states which are normally oxidising in aqueous solution can be stabilised by complexing e.g. compare the Co(II)/Co(III) potential when complexed with water (+1.82V) and with the ligand ammonia (+0.10V).

    • There are classic examples of disproportionation where an intermediate oxidation state species spontaneously changes into a higher and lower' oxidation state species e.g. the disproportionation reactions

      • Cu(I) ==> Cu(0) + Cu(II) and Mn(VI) ==> Mn(II) + Mn(VII).

      • These are described in detail, complete with electrode potential arguments for thermodynamic feasibility, under the respective metal.

    • How do you work out what will oxidise what? or what will reduce what?

      • How to work out the feasibility of reaction from electrode potential data is described in Appendix 5.

      • Using an electrode potential chart like the one above or a list of redox potentials the following rules apply.

        • To facilitate an oxidation, the half–cell potential of the oxidising agent must be less negative or more positive than the redox potential of the 'system' you wish to oxidise.

          • So using at the redox potential chart for example:

          • Dissolved oxygen will oxidise Co2+ to Co3+ in presence of ammonia – forms the amine complexes, but the hexaaqua complex of Co2+ is stable in the presence of oxygen if no ammonia present.

            • Co3+/Co2+ (H2O ligand, EØ = +1.82V), O2/H2O

            • (EØ = +1.23, less than +1.82 but more than +0.10V), Co3+/Co2+ (NH3 ligand, EØ = +0.10V)

            • So [Co(H2O)6]2+ is stable in the presence of oxygen, but [Co(NH3)6]2+ will be oxidised to [Co(NH3)6]3+.

            • You could then further predict that [Co(H2O)6]3+ will oxidise water to oxygen!

        • To facilitate a reduction, the half–cell potential of the reducing agent must be more negative or less positive than the redox potential of the 'system' you wish to reduce.

          • So using the redox potential chart and the half–cell redox potential for I2/I of +0.54V:

            • hexaaquairon(III) ions will be reduced by iodide ions because EØ for Fe3+/Fe2+ (H2O ligand) is +0.77V

            • i.e. the Fe3+ will oxidise the iodide ions rather than iodine oxidising the Fe2+ ions.

            • [Fe(H2O)6]3+ is reduced to [Fe(H2O)6]2+, iron(III) to iron(II).

            • However if the ligand is the cyanide ion, then iodide ions will not reduce the Fe3+ cyanide ion complex but iodine would oxidise [Fe(CN)6]4– to [Fe(CN)6]3–, iron(II) to iron(III).

            • Fe3+/Fe2+ (H2O ligand, EØ = +0.77V), I2/I

            • (EØ = +0.54, less than +0.77 but more than +0.36V), Fe3+/Fe2+ (CN ligand, EØ = +0.36V)

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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