10.7. Chemistry of Manganese Mn, Z=25, 1s²2s²2p⁶3s²3p⁶3d⁵4s²

data comparison of manganese with the other members of the 3d–block and transition metals

Extended data table for MANGANESE

• Uses of MANGANESE

  • Manganese is a hard brittle metal but is a most important metal in steels making tough and hard wearing manganese steel alloys e.g. for use as rock drills and railway points.

  • Ferromanganese (80% Mn, 20% Fe) is commonly used in such alloys.

  • Several compounds of manganese are used in the laboratory and manufactured products as oxidising agents.

    • Potassium manganate(VII), KMnO4, is used as a volumetric agent in redox analysis titrations.

    • Manganese(IV) oxide, MnO₂, is used in zinc–carbon batteries (oxidises hydrogen to water).

  • Manganese(IV) oxide, MnO₂, is also used in paints dyes.

  • Manganese compounds are found in fertilisers, fungicides and herbicides.

• Biological role of manganese

  • Manganese is an essential trace element.

  • Manganese activates the enzyme alkaline phosphatase in bone formation and the enzyme organase in urea formation.

  • Manganese deficiency can cause deformation of the skeleton and sterility.

  • In plants manganese ions activate carboxylases.

The Chemistry of MANGANESE

• Manganese(II) oxidation state chemistry

• The reactions of the manganese(II) ion:

  • An aqueous solution of manganese(II) sulfate MnSO_4(aq) or manganese(II) chloride MnCl_2(aq) will do for most laboratory experiments investigating the chemistry of the manganese(II) state.

  • Manganese(II) salts are readily made by dissolving the carbonate, MnCO₃, in the appropriate dilute acid.

    • e.g. MnCO_3(s) + 2HCl_(aq) ==> MnCl_2(aq) + H₂O_(l) + CO_2(g)

    • H₂SO₄ for the sulphate, MnSO₄,  and 2HNO₃ for the nitrate, Mn(NO₃)₂.

    • The very pale pink hexaaquamanganese(II) [Mn(H₂O)₆]²⁺ is quite 'redox' stable in aqueous solution.

  • From manganese(II) solutions, the alkalis sodium hydroxide or ammonia, produce the hydrated manganese(II) hydroxide precipitate. There is no further reaction with excess of either alkali.

    • Mn²⁺_(aq) + 2OH^–_(aq) ==> Mn(OH)_2(s)  (can be written as the neutral complex [Mn(OH)₂(H₂O)₄]

      • the hydroxide is almost white if oxygen excluded, but it gradually turns brown to form hydrated manganese(III) oxide.

      • then 4Mn(OH)_2(s)  + O_2(g) ==> 2Mn₂O_3(s)  + 4H₂O_(l)   

        • the hydrated oxide Mn₂O₃can also be written as a hydroxy–oxide, MnO(OH)

      • Mn oxidised (II)==>(III) and ==>(IV) possibly MnO₂?, O reduced (0)==>(–1)

    • VIEW ppts. with OH^–, NH₃ and CO₃^2–, & complexes with excess reagent

  • With manganese(II) ion solutions, alkaline aqueous sodium carbonate solutions produces a precipitate of manganese(II) carbonate.

    • Mn²⁺_(aq) + CO₃^2–_(aq) ==> MnCO_3(s) (white ppt.)

      • Like Mn(OH)₂ it readily oxidises to brown Mn₂O₃.

  • Oxidation of the manganese(II) ion

    • Acidified Mn²⁺ is not oxidised by hydrogen peroxide H₂O₂.

    • BUT alkaline Mn(OH)₂ + H₂O₂ gives brown Mn₂O₃ or MnO(OH), a hydrated manganese(III) oxide/hydroxide.

    • This again illustrates how redox potentials vary with pH i.e. change in relative stability of oxidation states for the Mn³⁺/Mn²⁺ half–cell potential.

  • The hexa–aqua manganese(II) ion readily forms complexes with polydentate ligands.

  • (i) [Mn(H₂O)₆]²⁺_(aq) + 3en_(aq) ===> [Mn(en)₃]²⁺_(aq) + 6H₂O_(l)   (en = H₂NCH₂CH₂NH₂)

    • K_stab = {[Mn(en)₃]²⁺_(aq)} / {[Mn(H₂O)₆]²⁺_(aq)} [en_(aq)]³

    • K_stab = 5.0 x 10⁵ mol^–3 dm⁹ [lg(K_stab) = 5.7]

  • (ii) [Mn(H₂O)₆]²⁺_(aq) + EDTA^4–_(aq) ===> [Mn(EDTA)]^2–_(aq) + 6H₂O_(l)

    • K_stab = {[Mn(EDTA)₃]^2–_(aq)} / {[Mn(H₂O)₆]²⁺_(aq)} [EDTA^4–_(aq)]

    • K_stab = 1.0 x 10¹⁴ mol^–1 dm³ [lg(K_stab) = 14.0]

  • The higher K_stab value for EDTA reflects the greater entropy change. A simplistic, but not illegitimate argument, shows that in (i) a net gain of 3 particles, but in (ii) 5 more particles are formed.

• The electrode potential chart highlights the values for various oxidation states of manganese.

• Summary of some complexes–compounds & oxidation states of manganese compared to other 3d–block elements

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• Manganese(III) oxidation state chemistry

• The chemistry of the manganese(III) ion

  • The violet Mn(H₂O)₆]³⁺(aq) ion is unstable in aqueous solution.

    • I don't think you need to know about this.

• Manganese(IV) oxidation state chemistry

• The chemistry of manganese(IV) in terms of manganese (IV) oxide

  • The only important manganese(IV) compound is the solid black oxide, MnO₂.

  • Manganese(IV) oxide is an excellent catalyst for the decomposition of hydrogen peroxide which is a useful way of making oxygen for school laboratory experiments. (See Gas Preparations)

    • 2H₂O_2(aq) ==> 2H₂O_(l) + O_2(g)

  • If a small quantity of manganese(IV) oxide is added to ice–cooled concentrated hydrochloric acid an anionic octahedral manganese(IV) chloro complex ion is formed. If the mixture is filtered through glass wool the brown colour of the complex can be seen.

    • (i) MnO_2(s) + 4H⁺_(aq) + 6Cl^–_(aq) ==> [MnCl₆]^2–_(aq) + 2H₂O_(l)

    • If the mixture is warmed, chlorine is formed as the complex decomposes to Mn(II) compounds.

      • (ii) [MnCl₆]^2–_(aq) ==> MnCl_2(aq) + 2Cl^–_(aq) + Cl_2(g)

    • The overall equation is (iii) MnO_2(s) + 4H⁺_(aq) + 4Cl^–_(aq) ==>MnCl_2(aq) + Cl_2(g) + 2H₂O_(l)

• The chemistry of manganese(VI) oxidation state

  • A solution of the tetrahedral dark green manganate(VI) ion, MnO₄^2– can be made by strongly heating a mixture of manganese(IV) oxide, potassium hydroxide and potassium chlorate(V) in a crucible and extracting the manganese(VI) compound with water.

    • 3MnO₂ + 6OH^– + ClO₃^– ==> 3MnO₄^2– + 3H₂O + Cl^–

    • See for full redox analysis of the reaction

  • However, the manganate(VI) ion, MnO₄^2– is unstable, especially in acid solution, and slowly undergoes disproportionation – i.e. a species in one oxidation state spontaneously and simultaneously changes into two species of different oxidation states – one higher and one lower in oxidation number. Adding dil. sulphuric acid to the crucible fusion extract will hasten the process.

  • The green solution of the manganate(VI) ion changes to the purple manganate(VII) ion and a black precipitate of manganese(IV) oxide is formed.

  • 3MnO₄^2–_(aq) + 4H⁺_(aq) ==> 2MnO₄^–_(aq) + MnO_2(s) + 2H₂O_(l)

  • The equilibrium constant for the reaction, K, is ~10⁵⁸, so there ain't much chance of the green colour hanging around after acidification!

  • The oxidation state changes are 3Mn(+6) ==> 2Mn(+7) + Mn(+4),

  • • a total '18 units worth' of redox change, always check your oxidation number balancing before anything else! The oxidation states of H and O remain at +1 and –2 respectively.

  • The obviously feasible and spontaneous disproportionation reaction can be explained by considering the standard electrode potentials (standard reduction potential) involved (quoted as half–cell reductions, as is the convention).

    • (i) MnO₄^–_(aq) + e^– ==> MnO₄^2–_(aq)  (E^Ø = +0.56V)

    • (ii) MnO₄^2–_(aq) + 4H⁺_(aq) + 2e^–  ==> MnO_2(s) + 2H₂O_(l)  (E^Ø = +1.70V, in acid solution)

    • (ii) has the more positive potential, so this will be the reduction half–cell reaction.

    • (i) has the less positive potential, so this will be (reversed) the oxidation half–cell reaction.

    • E^Ø_reaction = E^Ø_reduction – E^Ø_oxidation = (+1.70) – (0.56) = +1.14V, well over 0V, therefore very feasible!

    • Incidentally:

      • Given the two half–cell reactions, you get the complete balanced equation by adding (ii) plus 2 x (i) reversed.

      • The greater stability of the manganate(VI) ion in alkali can also be explained by considering the electrode potential for (ii) in an alkaline media.

      • (iii)  MnO₄^2–_(aq) + 2H₂O_(l) + 2e^–  ==> MnO_2(s) + 4OH^–_(aq)  (E^Ø = +0.59V, in alkaline solution)

      • so, re–calculating gives

      • E^Ø_reaction = E^Ø_reduction – E^Ø_oxidation = (+0.59) – (0.56) = +0.03V, just over 0, therefore just feasible! but on the basis of an equilibrium argument, here, the far lower E^Ø_reaction, suggests the MnO₄^2– ion is far more likely to exist, i.e. more stable, in a very high pH solution and in practice it is stable for a few hours in alkali.

      • This is a good example of how change in pH can affect a standard reduction potential.

      • See also copper(I) chemistry for another example of disproportionation.

• The chemistry of the manganese(VII) oxidation state i.e. the manganate(VII) ion

  • The tetrahedral deep purple manganate(VII) ion, MnO₄^–, can be considered as an intensely coloured and very stable complex ion (except in the presence of something that is readily oxidised!).

  • Potassium manganate(VII), KMnO₄ is used to titrate (i) iron(II) ions, (ii) ethanedioates, (iii) hydrogen peroxide and (iv) nitrate(III) ions (old name 'nitrite').

  • The titrations are done with dilute sulphuric acid present to prevent side reactions e.g. MnO₂ formation (brown colouration or black precipitate).

  • The mineral acid must be dilute sulfuric acid because potassium manganate(VII) will oxidise hydrochloric acid (Cl^– ==> Cl₂) and nitric(V) acid is an oxidising agent itself, so use of either of these acids leads to inaccurate false titration results.

  • The Mn²⁺ ions formed are almost colourless (very pale pink), so the end–point is the first permanent faint pink due to the first trace of excess of the brilliant purple manganate(VII) ion. 

  • (i) MnO₄^–_(aq) + 8H⁺_(aq)  + 5Fe²⁺_(aq) ==> Mn²⁺_(aq) + 5Fe³⁺_(aq) + 4H₂O_(l)

  • Theoretically, there are actually two simultaneous colour changes, both masked by the redox indicator change.

    • The purple manganate(VII) ion changes on reduction to the very pale pink manganese (II) ion,

    • and the pale green iron(II) ion changes on oxidation to the orange iron(III) ion,

    • However, in the dilute solution of the titration mixture, the first permanent pink colour does stand out from the pale orange of the iron(III) ion plus the very pale pink of the manganese(II) ion.

    • In the other examples (ii) to (iv) below, the reductants and oxidation products are colourless, so the colour of the very pale pink manganese(II) ion is visually overridden by the first drop of excess of the bright purple potassium manganate(VII).

  • (ii) 2MnO₄^–_(aq) + 16H⁺_(aq)  + 5C₂O₄^2–(aq) ==> 2Mn²⁺_(aq)  + 8H₂O_(l) + 10CO_2(g)

    • this reaction speeds up as the titration proceeds because Fe³⁺ acts as a catalyst, this situation is known as auto–catalysis.

  • (iii) 2MnO₄^–_(aq) + 6H⁺_(aq)  + 5H₂O_2(aq) ==> 2Mn²⁺_(aq) + 8H₂O_(l) + 5O_2(g)

  • (iv) 2MnO₄^–_(aq) + 6H⁺_(aq)  + 5NO₂^–_(aq) ==> Mn²⁺_(aq) + 5NO₃^–_(aq) +  3H₂O_(l)

  • See also fully worked examples of redox volumetric titration calculation questions.

  • The autocatalysis of the ethanedioate/potassium manganate (VII) titration reaction by the Mn²⁺ ions is described under homogeneous catalysis in Appendix 6.

  • Potassium manganate(VII), is strong enough to oxidise chloride ions. Running conc. hydrochloric acid onto the damp crystals is a handy way of making chlorine in the laboratory.

    • 2MnO₄^–_(aq) + 16H⁺_(aq)  + 10Cl^–_(aq) ==> 2Mn²⁺_(aq)  + 8H₂O_(l) + 5Cl_2(g)

    • This reaction is the reason that dilute sulphuric acid is used in potassium manganate(VII) titrations and not dil. hydrochloric acid, which would lead to inaccurate results.

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

keywords redox reactions ligand substitution displacement redox reactions ligand substitution displacement balanced equations formula complex ions complexes ligand exchange reactions redox reactions ligands colours oxidation states manganese ions Mn(0) Mn2+ Mn(+2) Mn(II) Mn3+ Mn(+3) Mn(III) Mn4+ Mn(+4) Mn(IV) Mn(+6) Mn (VI) Mn(+7) Mn(VII): MnSO4 MnCl2 MnO2 MnO Mn2O3 MnO4– MnO42– KMnO4 Mn(OH)2 MnCO3 + 2HCl ==> MnCl2 + H2O + CO2 Mn2+ + 2OH– ==> Mn(OH)2 [Mn(OH)2(H2O)4] 4Mn(OH)2 + O2 ==> 2Mn2O3 + 4H2O  Mn2+ + CO32– ==> MnCO3 [Mn(H2O)6]2+ + 3en  ===> [Mn(en) 3]2+ + 6H2O (en = H2NCH2CH2NH2) Kstab = {[Mn(en)3]2+} / {[Mn(H2O)6]2+} [en]3 Kstab = 5.0 x 105 mol–3 dm9 [lg(Kstab) = 5.7] [Mn(H2O)6]2+ + EDTA4– ===> [Mn(EDTA)]2– + 6H2O Kstab = {[Mn(EDTA)3]2–} / {[Mn(H2O)6]2+} [EDTA4–] MnO2 + 4H+ + 6Cl– ==> [MnCl6]2– + 2H2O [MnCl6]2– ==> MnCl2 + 2Cl– + Cl2 MnO2 + 4H+ + 4Cl–  ==>MnCl2 + Cl2 + 2H2O 3MnO2 + 6OH– + ClO3– ==> 3MnO42– + 3H2O + Cl– MnO42– + 4H+ ==> 2MnO4– + MnO2 + 2H2O 3Mn(+6) ==> 2Mn(+7) + Mn(+4) MnO4– + e– ==> MnO42– (EØ = +0.56V) MnO42– + 4H+ + 2e– ==> MnO2 + 2H2O MnO42– + 2H2O + 2e– ==> MnO2 + 4OH– MnO4– + 8H+ + 5Fe2+ ==> Mn2+ + 5Fe3+ + 4H2O 2MnO4– + 16H+  + 5C2O42– ==> 2Mn2+ + 8H2O + 10CO2 2MnO4– + 6H+ + 5H2O2 ==> 2Mn2+ + 8H2O + 5O2 2MnO4– + 6H+ + 5NO2– ==> Mn2+ + 5NO3– + 3H2O 2MnO4– + 16H+ + 10Cl– ==> 2Mn2+ + 8H2O + 5Cl2 oxidation states of manganese, redox reactions of manganese, ligand substitution displacement reactions of manganese, balanced equations of manganese chemistry, formula of manganese complex ions, shapes colours of manganese complexes  Na2CO3 NaOH NH3

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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, 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