Doc Brown's Chemistry  Periodic Table Revision Notes - Part 10. 3d block - Transition Metals

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

Revision notes for GCE Advanced Subsidiary Level AS Advanced Level A2 IB Revise AQA GCE Chemistry OCR GCE Chemistry Edexcel GCE Chemistry Salters Chemistry CIE Chemistry revising courses for pre-university students (equal to US grade 11 and grade 12 and Honours/honors level courses) GCSE Periodic Table notes * GCSE notes Transition Metals * EMAIL comment

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, E^ø * 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!

 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 FeO₄^2-, 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, E^Ø = +1.82V), O₂/H₂O (E^Ø = +1.23, less than +1.82 but more than +0.10V), Co³⁺/Co²⁺ (NH₃ ligand, E^Ø = +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 E^Ø 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, E^Ø = +0.77V), I₂/I^- (E^Ø = +0.54, less than +0.77 but more than +0.36V), Fe³⁺/Fe²⁺ (CN^- ligand, E^Ø = +0.36V)

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

Copyright © Dr W P Brown 2000-2010 All rights reserved on the revision notes pages, quizzes, worksheets, x-words etc.