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 3d block-transition metal oxidation states, electrode potentials of hydrated ions & other complex ions

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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 and electrode potentials for half-equations

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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

Doc Brown's Chemistry Advanced Level Pre-University Chemistry Revision Study Notes for UK IB KS5 A/AS GCE advanced level inorganic chemistry students US K12 grade 11 grade 12 inorganic chemistry - 3d block transition metal chemistry Sc Ti V Cr Mn Fe Co Ni Cu Zn

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!

Summary of oxidation states of the 3d block metals (least important) Ti to Cu are true transition metals

Sc Ti V Cr Mn Fe Co Ni Cu Zn
                +1  
  (+2) (+2) (+2) +2 +2 +2 +2 +2 +2
+3 +3 +3 +3 (+3) +3 +3 (+3) (+3)  
  +4 +4   +4     (+4)    
    +5              
      +6 (+6) (+6)        
        +7          
3d14s2 3d24s2 3d34s2 3d54s1 3d54s2 3d64s2 3d74s2 3d84s2 3d104s1 3d104s2
The outer electron configurations (beyond [Ar])

Examples are illustrated below.

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-cyan(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) NiAs

and various complexes

K3CuF6

contains the

[CuF6]3- ion

+4, (IV) TiO2 white(s)

[TiO]2+ colourless(aq)

TiCl4 colourless(l)

[VO]2+ blue(aq) MnO2 black(s) K2NiF6

contains the

[NiF6]2- ion

+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. 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.

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

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

    6. 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 and relative stability of oxidation states:

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.

Some trends on the nature and stability of transition metal ions:

The lower oxidation states of transition metals are usually found as simple ionic compounds e.g. containing ions such as Cr3+, Mn2+, Fe2+, Co2+, Ni2+, and Cu2+.

The transition metal compounds of their higher oxidation states are usually bound to an relatively highly electronegative element such as oxygen or fluorine in an anion

e.g. vanadate(V) VO4-, manganate(VII) MnO4- and ferrate(VI) FeO42-.

These are potentially strong oxidising agents e.g. potassium manganate(VII).

You do not usually get simple ions like V4+, Fe6+ or Mn7+.

For the 3d block in general in terms of redox reactions:

From left to right, the higher oxidation states become less stable relative to lower oxidation states.

Compounds of transition metals in a high oxidation state tend to be oxidising agents (comment above).

Compounds of transition metals in a low oxidation state tend to be reducing agents e.g. Ti2+ and Fe2+.

The relative stability of the +2 state relative to the +3 state increases from left to right e.g. the hydrated aqueous ions of Ti2+ and Fe2+.are easily oxidised to Ti3+ and Fe3+, but difficult to oxidise Co2+, Ni2+ and Cu2+ to Co3+, Ni3+ and Cu3+.

BUT, take care, complexing the central ion with ligands other than water and considerably change the relative stability of the complex i.e. can cause the half-cell potential to be significantly changed.

All except scandium (Sc/Sc3+), which is not that reactive towards 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 tend to be the 'simple' hydrated ions such as ...

Sc3+ (only ion), [TiO]2+, VO2+, Cr3+, Mn2+, Fe3+, Co2+, Ni2+, Cu2+ and Zn2+ (only ion), NOT Ti2+, Cr2+ or Fe2+.

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.

Other comments on the relative stability of transition metal ions

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?

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 ion of Co2+ is stable in the presence of oxygen if no ammonia present.

Co3+/Co2+ (H2O ligand, EØ = +1.82V), O2/H2O (EØ = +1.23V in neutral solution)

(EØ = +1.23 is 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 can 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)

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

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


INORGANIC Part 10 3d block TRANSITION METALS sub–index:

10.1–10.2 Introduction to 3d–block Transition Metal chemistry

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 of complexes and entropy changes

Appendix 9. Colorimetric analysis and complex ion formula * Appendix 10 3d block – extended data table

Appendix 11 3d–block compounds - transition metal 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

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