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

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!

Examples are illustrated below.

 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 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.

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) FeO₄^2-.

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 for the number of water ligands dative covalently bonded to the central metal ion..

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 for copper and manganese

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, E^Ø = +1.82V), O₂/H₂O (E^Ø = +1.23V in neutral solution)

(E^Ø = +1.23 is 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 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 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)

(E^Ø = +0.54, less than +0.77 but more than +0.36V), Fe³⁺/Fe²⁺ (CN^– ligand, E^Ø = +0.36V)

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