Scheme of chemical properties of manganese

Low oxidation states

Manganese forms a decacarbonyl Mn₂(CO)₁₀ in which each manganese has the required share in 18 electrons to achieve the noble gas configuration. Reduction of this covalent compound with sodium amalgam gives the salt Na[Mn(CO)₅], sodium pentacarbonylmanganate (-1); in the ion Mn(CO)₅ the noble gas structure is again attained.

Oxidation state + 2

This is the most common and stable state of manganese; the five d electrons half fill the five d-orbitals, and hence any transition of d electrons is not favourable in energy. The colour is correspondingly pale (usually pink).

In oxidation state +2, manganese shows fewer 'transition-like' characteristics than any other transition metal ion (d⁵ configuration); thus the aquo-ion [Mn(H₂O)₆]²⁺ forms a “normal” carbonate MnCO₃ which is insoluble in water and occurs naturally as “manganese spar”. The aquo-ion forms typical hydrated salts, for example MnSO₄∙7H₂O, MnCl₂∙ x H₂O and double salts, for example (NH₄)₂Mn(SO₄)₂∙6H₂O.

Dehydration of the simple hydrated salts, by heating, produces the anhydrous salt without decomposition. Addition of alkali precipitates the white basic manganese(II) hydroxide Mn(OH)₂.

Addition of ammonia to an aqueous solution of a manganese(II) salt precipitates Mn(OH)₂; reaction of ammonia with anhydrous manganese(II) salts can yield the ion [Mn(NH₃)₆]²⁺.

MnO. Preparation: Decomposition of manganese salts(MnCO₃, Mn(NO₃)₂) and MnO₂ in hydrogen atmosphere:

MnO₂ + H₂ = MnO + H₂O

Reducing properties Mn(ІІ)

All Mn(ІІ) compounds display reducing properties

When annealing MnO forms to Mn₃O₄ (Mn⁺²Mn⁺³₂O₄) at the presence of air:

6MnO + O₂ = 2Mn₃O₄.

Mn(OH)₂ is oxidised by air at standard conditions and becomes dark:

2Mn(OH)₂ + O₂ + 2H₂O = 2Mn(OH)₄ ®MnO(OH)₂ + H₂O

Mn(OH)₄ has a various composition (MnO₂)_x(H₂O)_y. The oxidation process has the following steps:

4Mn(OH)₂ + O₂ + 2H₂O = 4Mn(OH)₃

4Mn(OH)₃ + O₂ + 2H₂O = 4Mn(OH)₄

Mn(OH)₂ react with halogens (Cl₂, Br₂) in alkaline medium:

Mn(OH)₂ + Br₂ + 2NaOH = MnO₂ + 2NaBr + 2H₂O.

It also transforms to Mn(VI) in concentrated alkali solution (t^o_, a plenty of oxidant):

3MnSO₄ + 2KClO₃ + 12KOH = 3K₂MnO₄ + 2KCl + 3K₂SO₄ + 6H₂O

Mn(ІІ) compounds are stable in neutral and acid mediums to the action of Cl₂, Br₂ and O₂. The stronger oxidants can oxidise Mn(ІІ), for instance, ozone:

MnSO₄ + O₃ + 3H₂O = Mn(OH)₄ + H₂SO₄ + O₂.

It is a qualitative reaction on the presence of ozone in gas mixtures.

Mn(II) can be oxidised to Mn(VII) in acid medium when it reacts with very strong oxidants:

2MnSO₄ + 5PbO₂ + 6HNO₃ = 2HMnO₄ + 2PbSO₄ + 3Pb(NO₃)₂ + 2H₂O;

2MnSO₄ + 5(NH₄)₂S₂O₈ + 8H₂O = 2HMnO₄ + 5(NH₄)₂SO₄ + 7H₂SO₄.

dation state + 3

Сompounds Mn(ІІІ) and Re(ІІІ) are of low stability, especially in aqueous solutions. Mn (III) is reduced easily to Mn(II) or disproportionate forming Mn(II) and Mn(IV). Re (ІІІ), vice versa, is a reductant that forms highest oxidation states.

Оxide Mn(ІІІ) is produced by heating of MnO₂:

4MnO₂ = 2Mn₂O₃ + O₂.

Mn₂O₃ is a basic oxide:

Mn₂O₃ + 3H₂SO₄ = Mn₂(SO₄)₃ + 3H₂O

Mn(OH)₃ can be prepared by the reaction:

Mn₂(SO₄)₃ + 6KOH = 2Mn(OH)₃ + 3K₂SO₄

Hydrated Mn(ІІІ) oxide (or to simplify - Mn(OH)₃) is an intermediate product of Mn(OH)₂ oxidation by oxygen. Mn(OH)₃ displays reducing properties in alkaline medium:

2Mn(OH)₃ + Br₂+ 2KOH = 2MnO₂ + 2KBr + 4H₂O

Oxidation state + 3

This state is unstable with respect to disproportionation in aqueous solution:

2Mn³⁺(aq) + 2H₂O -> Mn²⁺ (aq) + MnO₂ + 4H⁺.

However the Mn³⁺(aq) ion can be stabilised by using acid solutions or by complex formation; it can be prepared by electrolytic oxidation of manganese(II) solutions. The alum CaMn(SO₄)₂^.12H₂O contains the hydrated Mn³⁺ ion, which is strongly acidic. The complexes of manganese(III) include [Mn(CN)₆]^3- (formed when manganese(II) salts are oxidised in presence of cyanide ions), and [MnF₅(H₂O)]^2-, formed when a manganese(II) salt is oxidised by a manganate(VII) in presence of hydrofluoric acid:

4Mn²⁺ + 8H⁺ + MnO₄ = 5Mn³⁺ + 4H₂O

Mn³⁺ + H₂O + 5F^- = [MnF₅(H₂O)]^2-

Oxidation of manganese(II) hydroxide by air gives the brown hydrated oxide Mn₂O₃aq, and this on drying gives MnO(OH) which occurs in nature as manganite. (The oxide Mn₂O₃ also occurs naturally as braunite.) Heating of the oxide Mn₂O₃ gives the mixed oxide Mn₃O₄ [manganese(II) dimanganese(III) oxide].

In general, manganese(III) compounds are coloured, and the complexes are octahedral in shape; with four d electrons, the colour is attributable in part to d-d transitions.

Mn₂O₃ is a solid black substance that dissolves slightly in acids and water. Preparation: Decomposition of manganese MnO₂:

4MnO₂ = 2Mn₂O₃ + O₂

compounds(IV)

The oxidation state +4 is the relatively stable state of Mn, Tc and Re. Compounds of Mn(IV) are not diverse due to stability and low solubility of MnO₂ . Namely MnO₂ is the most typical product of oxidation of Mn(II) and reduction Mn(VII), Mn(VI) in neutral medium. The other compounds Mn(IV) are not stable and can be easily transformed to MnO₂.

MnO₂ is abundant in nature. Manganese(IV) oxide is a dark-brown solid, insoluble in water and dilute acids. Its catalytic decomposition of potassium chlorate(V) and hydrogen peroxide has already been mentioned. It is known that manganese dioxide form family of tunnelled polymorphs capable of intergrowth of their tunnels. Manganese dioxide exists in five polymorphic modifications: a-, b-, g-, e-, d-, and l-MnO₂ (see below).

Pure MnO₂ is obtained by strong heating of manganese(II) nitrate:

Mn(NO₃)₂ = MnO₂ + 2NO₂,

MnO₂ is an amphoteric oxide although its acid and basic properties are expressed weakly. It forms low stable sulfate with concentrated H₂SO₄:

MnO₂ + 2H₂SO₄ = Mn(SO₄)₂ + 2H₂O

And unstable salts Mn(IV) manganites with alkalis:

MnO₂ + 2KOH = K₂MnO₃ + H₂O metamanganite

MnO₂ + 4KOH = K₄MnO₄ + 2H₂O orthomanganite

Hausmanite Mn₃O₄ can be represented as Mn(II) orthomanganite:

2Mn(OH)₂ + H₄MnO₄ = Mn₂MnO₄ + 4H₂O

Redox properties of Mn(IV) compounds are typical. MnO₂ is a strong oxidant (E^o =1.23 V). The reaction with HCl illustrates this behaviour:

MnO₂ + 4HCl = MnCl₂ + Cl₂ + 2H₂O;

Although it dissolves in ice-cold concentrated hydrochloric acid forming the complex octahedral hexachloromanganate(IV) ion:

MnO₂ + 6HCl = [Mn^IVCl₆]^2- + 2H⁺ + 2H₂O

2MnO₂ + 2H₂SO₄ = 2MnSO₄ + O₂ + 2H₂O.

When melting with alkalis at the presence of oxidants it displays reducing agent properties:

2MnO₂ + 4KOH + O₂ = 2K₂MnO₄ + 2H₂O

MnO₂ + 2KOH + KNO₃ K₂MnO₄ + KNO₂ + H₂O

2MnO₂ + 3PbO₂ + 6HNO₃ = 2HMnO₄ + 3Pb(NO₃)₂ + 2H₂O

TcO₂ (black) and ReO₂ (brown) oxides are known. ReO₂ can be obtained indirectly:

Re₂O₇ + 3H₂ = 2ReO₂ + 3H₂O

2NH₄ReO₄ = 2ReO₂ + N₂ + 4H₂O.

or when heated a mixture of Re₂O₇ and Re, disproportionation reaction:

2Re₂O₇ + 3Re = 7ReO₂

3ReO₃ = ReO₂ + Re₂O₇

Technetium(IV) oxide has similar preparation procedures. ReO₂ is unstable and disproportionate easily:

7ReO₃ = 3Re + 2Re₂O₇

Acid properties are decreased in the series Mn-Tc-Re. Re(OH)₄ displays very weak basic properties:

ReO₂ + 2NaOH = Na₂ReO₃ + H₂O renite

ReO₂ has reducing properties:

4ReO₂ +3O₂ = 2Re₂O₇

Oxidation state + 4. Analysis and complexes

An oxidation which can be used to estimate the amount of manganese(IV) oxide in a sample of pyrolusite is that of ethanedioicor oxalic acid, H₂C₂O₄:

MnO₂ + (COOH)₂ + H₂SO₄ = MnSO₄ + 2CO₂ + 2H₂O

Excess standard acid is added, and the excess (after disappearance of the solid oxide) is estimated by titration with standard potassium manganate(VII). Alternatively, a known weight of the pyrolusite may be heated with concentrated hydrochloric acid and the chlorine evolved passed into potassium iodide solution. The iodine liberated is titrated with sodium thiosulfate:

MnO₂ = Cl₂= I₂ = 2S₂O₃^2-

Although the complex ion [MnCl₆]^2- is unstable, salts such as K₂[MnF₆] (containing the octahedral hexafluoromanganate(IV) ion) are much more stable and can be crystallised from solution.

Oxidation state + 5

This state exists as a manganate (V), the blue MnO₄^3-, Na₃MnO₄^.10H₂O is isomorphous with Na₃VO₄. When melting MnO₂ with soda and saltpeter аt the presence of air a typical blue colour of MnO₄^3- ion appears. Mn here has oxidation state +5:

MnO₂ + Na₂CO₃ + NaNO₃ = Na₃MnO₄ + CO₂ + NO₂

Salts Mn(+5) can easily disproportionate:

2Na₃MnO₄ + 2H₂O = MnO₂ + Na₂MnO₄ + 4NaOH

Oxidation state +6

This is only found in the green manganate(VI) ion. It is only stable in alkaline conditions; in neutral or acid solution itdisproportionates:

3 MnO₄^2- + 2H₂O = MnO₂ + 2MnO₄^- + 4OH^-

Manganates are stable in the strong alkaline medium. They have tendency to dissociate and form more stable compounds of Mn(IV I,VII):

3K₂MnO₄ + 2H₂O Û 2KMnO₄ + MnO₂ + 4KOH.

This reaction is reversible, the growth of alkali concentration stabilises MnO₄^2-, and vice versa, in acid medium Mn(VI) is decomposed completely:

3K₂MnO₄ + 2H₂SO₄ = 2KMnO₄ + MnO₂ + 2K₂SO₄ + 2H₂O.

It explains the fact that H₂MnO₄ has not been prepared yet and its anhydride MnO₃ does not exist.

Manganates are reducing agents when react with strong oxidants:

2K₂MnO₄ + Cl₂ = 2KMnO₄ + 2KCl.

There are more stable compounds of Tc(VI) and Re(VI). Oxides TcO₃ (red), ReO₃ (red), fluorides, chlorides, salts (for instance, K₂ReO₄) are known.

Compounds of Re(VI) disprportionate like Mn(VI).

ReO₃ is usually obtained:

ReO₂ + Re₂O₇ = 3ReO₃

Oxidation state + 7

Apart from two unstable oxyhalides, MnO₃F and MnO₃Cl, this state is exclusively represented by the oxide Mn₂O₇ and the anion MnO₄^-. Crystalline KMnO₄ forms oily deep-green or dark coloured liquid of Mn(VII) oxide with concentrated H₂SO₄:

2KMnO₄ + 2H₂SO₄ = Mn₂O₇ + 2KHSO₄ + H₂O.

This substance is very unstable. It is decomposed (sometimes explosively) on heating to manganese (IV) oxide and oxygen:

2Mn₂O₇ = 4MnO₂ + 3O₂.

Mn₂O₇ is a powerful oxidant. Reactons with organic substances (alcohols, ethers etc.) are so violent that organic compounds burn. Manganese heptoxide is an acid oxide. It forms with water permanganic acid.

Mn₂O₇ + H₂O = 2HMnO₄.

Dihydrate HMnO₄^.2H₂O can be isolated as purple solids by low temperature evaporation of the frozen solution. Permanganic acid is one of the most strong acids that dissociates into ions almost completely and has pink colour. This acid is stable in aqueous solutions (less than 20% by mass). It is actively decomposed in more concentrated solutions:

4HMnO₄ = 4MnO₂ + 3O₂ + 2H₂O.

Permanganic acid is also a violent oxidising agent, especially when it reacts with any organic material; it decomposes quickly at 276 K. The purple permanganate anion, MnO₄^- is tetrahedral; it owes its intense colour to charge transfer (since Mn⁺⁷ has no d-electrons). The potassium salt KMnO₄ is the most frequently used reagent but other cations form soluble permanganates too (all these salts are purple). Salts of large size cations (for example Cs⁺) are less soluble. Potassium permanganate can be prepared by electrolytic oxidation of manganese metal (oxidation from 0 to +7) using a manganese anode in potassium carbonate solution;

oxidation of manganate(II) (oxidation +2 to +7), using the peroxodisulfate ion S₂O₈^2-and a manganese(II) salt

Potassium manganate(VII) disproportionates on heating:

2KMnO₄ = K₂MnO₄ + MnO₂ + O₂

The manganate(VII) ion slowly oxidises water:

4 MnO₄^- + 4H⁺ = 4MnO₂ + 2H₂O + 3O₂

This reaction proceeds very slowly in absence of light, and aqueous solutions of potassium manganate(VII) are effectively stable for long periods if kept in dark bottles.

The manganate(VII) ion is one of the more useful oxidising agents; in acid solution we have

MnO₄^-(aq) + 8H₃O+ + 5e- = Mn²⁺(aq) + 12H₂O; E_o = + 1.52 V

Hence acid solutions of manganate(VII) are used in chemical analysis or to oxidise, for example, quantitatively. The equivalence point of redox titration is recognised by persistence of the purple colour. (Sulfuric acid is used to acidify, since hydrochloric acid is oxidised to chlorine, and nitric acid is an oxidising agent.)

Manganates(VII) are also applied extensively in organic chemistry, for example, to oxidise alcohols to aldehydes in acid or (more commonly) in neutral medium where manganese(IV) oxide is the product.

In concentrated alkali, manganese(VI) is more stable than manganese(VII) and the following reaction occurs:

4MnO₄^- + 4OH^- = 4MnO₄^2-+ 2H₂O + O₂

It can be seen that redox behaviour of permanganate-ion depends on pH index. The scheme illustrating this dependence is shown below:

2KMnO₄ + 5K₂SO₃ + 3H₂SO₄ = 2MnSO₄ + 6K₂SO₄ + 3H₂O

2KMnO₄ + 3K₂SO₃ + H₂O = 2MnO₂ + 3K₂SO₄ + 2KOH

2KMnO₄ + K₂SO₃ + 2KOH = 2K₂MnO₄ + K₂SO₄ + H₂O.

The relation of permanganates to halide-ions (Cl^-, Br^-, I^-) illustrates the decrease of their oxidizing abililty at the growth of pH index. MnO₄^- ion is able to oxidize all halides in an acid medium including chlorine:

2KMnO₄ + 16HCl = 2MnCl₂ + 5Cl₂ + 2KCl + 8H₂O.

It can oxidise only iodine in neutral medium. In contrary, Cl₂ and Br₂ are able to oxidize lower valenced manganese to +7:

2MnO₄^2- + Br₂ = MnO₄^- + 2Br^-.

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