* Advanced Inorganic Chemistry Transition metals 10.7 Manganese Chemistry Doc Brown's

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 Doc Brown's Chemistry  Periodic Table Revision Notes 10.7

Part 10. Transition Metals 3d-block:  10.7 Manganese Chemistry

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 * GCSE notes Transition Metals

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

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

  • Mn data table 1 summary * extended manganese data table 2 * Manganese & electrode potential chart 3d-block

  • The reactions of the manganese(II) ion:

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

      • e.g. MnCO3(s) + 2HCl(aq) ==> MnCl2(aq) + H2O(l) + CO2(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.

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

      • 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)  + O2(g) ==> 2Mn2O3(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 CO32-, & complexes with excess reagent

    • Alkaline aqueous sodium carbonate solutions produces a precipitate of manganese(II) carbonate.

      • Mn2+(aq) + CO32-(aq) ==> MnCO3(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)

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

      • Kstab = 5.0 x 105 mol-3 dm9 [lg(Kstab) = 5.7]

    • (ii) [Mn(H2O)6]2+(aq) + EDTA4-(aq) ===> [Mn(EDTA)]2-(aq) + 6H2O(l)

      • Kstab = {[Mn(EDTA)3]2-(aq)} / {[Mn(H2O)6]2+(aq)} [EDTA4-(aq)]

      • Kstab = 1.0 x 1014 mol-1 dm3 [lg(Kstab) = 14.0]

    • The higher Kstab 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 chemistry of manganese(III)

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

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

  • The chemistry of manganese(IV)

    • 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)

      • 2H2O2(aq) ==> 2H2O(l) + O2(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.

      • MnO2(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.

        • [MnCl6]2-(aq) ==> MnCl2(aq) + 2Cl-(aq) + Cl2(g)

  • The chemistry of manganese(VI)

    • A solution of the dark green manganate(VI) ion, MnO42- 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.

    • However, the manganate(VI) ion, MnO42- 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 changes to purple and a black precipitate is formed.

      • 3MnO42-(aq) + 4H+(aq) ==> 2MnO4-(aq) + MnO2(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- ==> MnO42-(aq)  (EØ = +0.56V)

        • (ii) MnO42-(aq) + 4H+(aq) + 2e-  ==> MnO2(s) + 2H2O(l)  (EØ = +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.

        • EØreaction = EØreduction - EØoxidation = (+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)  MnO42-(aq) + 2H2O(l) + 2e-  ==> MnO2(s) + 4OH-(aq)  (EØ = +0.59V, in alkaline solution)

          • so, re-calculating gives

          • EØreaction = EØreduction - EØoxidation = (+0.59) - (0.56) = +0.03V, just over 0, therefore just feasible! but on the basis of an equilibrium argument, here, the far lower EØreaction, suggests the MnO42- 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 manganese(VII):

    • The 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 Mn2+ ions formed are almost colourless, so the end-point is the first permanent faint pink due to the first trace of excess manganate(VII) ion.  The mineral acid must be dilute sulfuric acid because potassium manganate(VII) will oxidise hydrochloric acid (Cl- ==> Cl2) and nitric acid is an oxidising agent itself, so use of either acid leads to inaccurate results.

      • (i) MnO4-(aq) + 8H+(aq)  + 5Fe2+(aq) ==> Mn2+(aq) + 5Fe3+(aq) + 4H2O(l)

      • (ii) 2MnO4-(aq) + 16H+(aq)  + 5C2O42-(aq) ==> 2Mn2+(aq)  + 8H2O(l) + 10CO2(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)  + 5H2O2(aq) ==> 2Mn2+(aq) + 8H2O(l) + 5O2(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) + 5Cl2(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

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