Manganese transition metal Chemistry manganese(II) Mn2+ complex ions MnO4- manganate(VII) ‘permanganate’ manganese(IV) oxide MnO2 redox chemical reactions principal oxidation states +2 +4 +6 +7 GCE AS

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

Manganese exhibits
oxidation states of +2, +3, +4, +6 and +7, though the chemistry you will
most likely encounter is that of Mn2+ (+2) salts and complex ions, manganese(IV)
oxide, MnO2 (+4) and the useful oxidising agent (potassium)
manganate(VII) ion MnO4– (+7).

This page describes the
chemistry of the principal 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 and colours of
manganese complexes, formula of compounds

See also the
absorption
spectra and colours of manganese compounds *
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data comparison of manganese
with the other members of the 3d–block and transition metals

Elect.
pot. = standard electrode potential data for manganese
(EØ at 298K/25oC, 101kPa/1 atm.)

na = data not applicable to
manganese

Extended data table for MANGANESE

The oxidation states and electron
configuration of manganese
in the context of the 3d block of elements

The
electrode potential chart highlights the values for various
oxidation states of manganese +2 +3 +7.

Manganese oxidations states of +4 and +6 are
also mentioned in the text below.

The electrode potentials involving manganese
ions correspond to hydrated complex ions where the ligands are
water, oxide or hydroxide.

As you can see from the chart, changing either
the ligand or the oxidation state, will also change the
electrode potential for that half-reaction involving a manganese
ion.

The manganate(VII) ion is a powerful oxidising
agent.

A quick illustration of manganese oxidation
states

I came across an experiment on twitter (@mrspotassium)
where you stir an acidified potassium manganate(VII)
solution (in a beaker of conical flask) with sugary lollipop – mainly glucose sugar.

Alternatively, you can stir the solution
with a magnetic stirrer and carefully suspend the lollipop
into the potassium manganate(VII) solution – you can dip the
lollipop to a greater or lesser depth to control the rate of
reaction (a surface area rate factor!).

Glucose is a reducing agent and the
following reduction sequence occurs.

Quite
clever, the hard sugary lollipop only dissolves slowly,
slowing down the reactions so that you can see a series of
colour changes corresponding to the change in oxidation
state of manganese.

purple
Mn(VII) => green Mn(VI) => ? Mn(IV) => violet Mn(III) =>
pale pink Mn(II)

Not sure on the Mn(IV) colour, apart from
insoluble black solid MnO2, not sure if it can be
stabilised in solution? possibly with a suitable ligand?

More details via the
sub-index for this
page.

(b) Manganese(II) oxidation state chemistry

• The
  reactions of the manganese(II) ion:

  • Electron configuration of Mn2+
    is [Ar]3d5

  • An
    aqueous solution of manganese(II) sulfate MnSO4(aq)
    or manganese(II) chloride MnCl2(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.
      MnCO3(s)
      + 2HCl(aq) ===> MnCl2(aq) + H2O(l) +
      CO2(g)

    • H2SO4
      for the sulfate, 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]0

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

      • then
        4Mn(OH)2(s) + O2(g)
        ===> 2Mn2O3(s)
        + 4H2O(l)

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

    • VIEW more on ppts. with OH–, NH3
      and CO32–, & complexes with excess reagent

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

  • 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]

    • Remember [H2O] is not included in the
      equilibrium expression.

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

  • 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
  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

(c) Manganese(III) oxidation state chemistry

(d) Manganese(IV) oxidation state chemistry

(e) The
chemistry of manganese(VI) oxidation state

• A solution of the
  tetrahedral
  (O-Mn-O bond angle 109.5o)
  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. sulfuric 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.

• 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)

(f) 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 sulfuric
  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 bright purple potassium manganate(VII) at the end-point of
    the volumetric titration.

  • You do need excess dil. sulfuric acid
    in the titration and it will NOT act as an oxidising agent or a reducing
    agent to interfere with the quantitative and accurate volumetric redox
    titration.

  • e.g. acids you should NOT use
    for a potassium manganate(VII) redox titration:

    • conc. sulfuric acid is an oxidising
      agent, dilute is fine,

    • hydrochloric acid, manganate(VII) ion
      oxidises the chloride ion to chlorine,

    • nitric acid is an oxidising agent,

    • ethanoic acid is too weak to provide
      a high H+(aq) concentration and many other weak
      organic acids are oxidised

• (ii)
  2MnO4–(aq)
  + 16H+(aq) + 5C2O42–(aq)
  ===> 2Mn2+(aq) + 8H2O(l) +
  10CO2(g)

• (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 on these titrations.

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

biological role of manganese,
chemistry of the manganese(II) ion Mn2+, octahedral complexes of manganese(II),
standard electrode potential of Mn2+, oxidation of manganese(II) to
manganese(VII) with alkaline chlorine, chemistry of manganate(VII) ion MnO4 2-,
chemistry of the manganate(VI) ion,
hexaaquamanganese(II) ion, standard electrode potential for Mn2+
manganese(II) ion, oxidising power of manganate(VII), colours of manganese ions 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 transition metal manganese
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