Manganese(II) oxide- MnO - lower manganese oxide, monoxide.
basic oxide. Let's not dissolve in water. Easily oxidized to form a brittle MnO 2 shell. It is reduced to manganese when heated with hydrogen or active metals.
Manganese(II) oxide can be obtained by calcining at a temperature of 300 °C oxygen-containing salts of manganese(II) in an inert gas atmosphere. From common MnO 2 it is obtained through partial reduction at temperatures of 700-900 ° C with hydrogen or carbon monoxide.
Manganese(II) hydroxide- inorganic compound, manganese metal hydroxide with the formula Mn(OH) 2 , light pink crystals, insoluble in water. Shows weak basic properties. Oxidizes in air.
Manganese (II) hydroxide is formed by the interaction of its salts with alkalis:
Chemical properties.
Manganese (II) hydroxide is easily oxidized in air to brown manganese oxohydroxide, which further decomposes into manganese (IV) oxide:
· Manganese (II) hydroxide has basic properties. It reacts with acids and acid oxides:
· Manganese (II) hydroxide has reducing properties. In the presence of strong oxidizing agents, it can oxidize to permanganate:
Manganese(III) oxide- inorganic compound, manganese metal oxide with the formula Mn 2 O 3, brown-black crystals, insoluble in water.
Receipt.
· In nature, there are minerals brownite, kurnakite and bixbyite - manganese oxide with various impurities.
Oxidation of manganese(II) oxide:
Recovery of manganese(IV) oxide:
Chemical properties.
Decomposes on heating:
When dissolved in acids, it disproportionates:
When fused with metal oxides, it forms salts of manganites:
Does not dissolve in water.
Manganese(III) hydroxide –Mn2O3ּ H 2 O or MnO(OH) occurs naturally as a mineral manganite(brown manganese ore). Artificially obtained manganese (III) hydroxide is used as a black-brown paint.
When interacting with acidic oxidizing agents, it forms manganese salts.
Salts of manganese (II), as a rule, are well soluble in water, except for Mn 3 (PO 4) 2, MnS, MnCO 3.
manganese sulfate(II) MnSO 4 is a white salt, one of the most stable compounds of manganese (II). In the form of crystalline MnSO 4 7H 2 O occurs in nature. It is used for dyeing fabrics, and also, along with manganese (II) chloride MnCl 2 - to obtain other manganese compounds.
manganese carbonate(II) MnCO 3 is found in nature as manganese powder and is used in metallurgy.
manganese nitrate(II) Mn(NO 3) 2 is obtained only artificially and is used to separate rare earth metals.
Salts of manganese are catalysts for oxidative processes involving oxygen. They are used in desiccants. Linseed oil with the addition of such a desiccant is called drying oil.
Manganese(IV) oxide (manganese dioxide) MnO 2 - dark brown powder, insoluble in water. The most stable compound of manganese, widely distributed in the earth's crust (mineral pyrolusite).
Chemical properties.
Under normal conditions, it behaves rather inertly. When heated with acids, it exhibits oxidizing properties, for example, it oxidizes concentrated hydrochloric acid to chlorine:
With sulfuric and nitric acids, MnO 2 decomposes with the release of oxygen:
When interacting with strong oxidizing agents, manganese dioxide is oxidized to compounds Mn 7+ and Mn 6+:
Manganese dioxide exhibits amphoteric properties. So, when a sulfuric acid solution of the MnSO 4 salt is oxidized with potassium permanganate in the presence of sulfuric acid, a black precipitate of the Mn(SO 4) 2 salt is formed.
And when fused with alkalis and basic oxides, MnO 2 acts as an acid oxide, forming salts - manganites:
It is a catalyst for the decomposition of hydrogen peroxide:
Receipt.
Under laboratory conditions, it is obtained by thermal decomposition of potassium permanganate:
It can also be obtained by the reaction of potassium permanganate with hydrogen peroxide. In practice, the formed MnO 2 catalytically decomposes hydrogen peroxide, as a result of which the reaction does not proceed to the end.
At temperatures above 100 °C by reduction of potassium permanganate with hydrogen:
64. Manganese (VI) compounds, methods of preparation and properties. Manganese oxide (VII), permanganic acid and permanganates - obtaining, properties, application.
Manganese(VI) oxide- an inorganic compound, manganese metal oxide with the formula MnO 3, a dark red amorphous substance, reacts with water.
It is formed during the condensation of violet vapors released when a solution of potassium permanganate in sulfuric acid is heated:
Chemical properties.
Decomposes on heating:
Reacts with water:
Forms salts with alkalis - manganates:
Manganese(VI) hydroxide exhibits an acidic character. free manganese (VI) acid is unstable and disproportionates in an aqueous solution according to the scheme:
3H 2 MnO 4(c) → 2HMnO 4(c) + MnO 2(t) + 2H 2 O (l).
Manganates (VI) are formed by fusing manganese dioxide with alkali in the presence of oxidizing agents and have an emerald green color. Manganates (VI) are rather stable in strongly alkaline medium. When alkaline solutions are diluted, hydrolysis occurs, accompanied by disproportionation:
3K 2 MnO 4 (c) + 2H 2 O (l) → 2KMnO 4 (c) + MnO 2 (t) + 4KOH (c).
Manganates (VI) are strong oxidizing agents that are reduced in an acidic environment to Mn(II), and in neutral and alkaline environments - up to MNO2. Under the action of strong oxidizing agents, manganates (VI) can be oxidized to Mn(VII):
2K 2 MnO 4 (c) + Cl 2 (d) → 2KMnO 4 (c) + 2KCl (c).
When heated above 500 ° C, manganate (VI) decomposes into products:
manganate (IV) and oxygen:
2K 2 MnO 4 (t) → K 2 MnO 3 (t) + O 2 (g).
Manganese(VII) oxide Mn 2 O 7- greenish-brown oily liquid (t pl \u003d 5.9 ° C), unstable at room temperature; a strong oxidizing agent, in contact with combustible substances, ignites them, possibly with an explosion. Explodes from a push, from a bright flash of light, when interacting with organic substances. Manganese (VII) oxide Mn 2 O 7 can be obtained by the action of concentrated sulfuric acid on potassium permanganate:
The resulting manganese(VII) oxide is unstable and decomposes into manganese(IV) oxide and oxygen:
At the same time, ozone is released:
Manganese(VII) oxide reacts with water to form permanganic acid, which has a purple-red color:
It was not possible to obtain anhydrous permanganic acid; it is stable in solution up to a concentration of 20%. This is very strong acid, the apparent degree of dissociation in a solution with a concentration of 0.1 mol / dm 3 is 93%.
Permanganic acid – strong oxidizing agent . More energetic interaction Mn2O7 combustible substances ignite when in contact with it.
Salts of permanganic acid are called permanganates . The most important of these is potassium permanganate, which is a very strong oxidizing agent. Its oxidizing properties with respect to organic and inorganic substances are often encountered in chemical practice.
The degree of reduction of permanganate ion depends on the nature of the medium:
1) acidic environment Mn(II) (salts Mn 2+)
MnO 4 - + 8H + + 5ē \u003d Mn 2+ + 4H 2 O, E 0 \u003d +1.51 B
2) neutral environment Mn(IV) (manganese(IV) oxide)
MnO 4 - + 2H 2 O + 3ē \u003d MnO 2 + 4OH -, E 0 \u003d +1.23 B
3) alkaline environment Mn (VI) (manganates M 2 MnO 4)
MnO 4 - +ē \u003d MnO 4 2-, E 0 \u003d + 0.56B
As can be seen, the strongest oxidizing properties of permanganates are exhibited by in an acidic environment.
The formation of manganates occurs in a highly alkaline solution, which suppresses hydrolysis K2MnO4. Since the reaction usually takes place in sufficiently dilute solutions, the end product of the reduction of permanganate in an alkaline medium, as well as in a neutral one, is MnO 2 (see disproportionation).
At a temperature of about 250 ° C, potassium permanganate decomposes according to the scheme:
2KMnO 4(t) K 2 MnO 4(t) + MnO 2(t) + O 2(g)
Potassium permanganate is used as an antiseptic. Aqueous solutions of its various concentrations from 0.01 to 0.5% are used for wound disinfection, gargling and other anti-inflammatory procedures. Successfully 2 - 5% solutions of potassium permanganate are used for skin burns (the skin dries up, and the bubble does not form). For living organisms, permanganates are poisons (cause proteins to coagulate). Their neutralization is carried out with a 3% solution H 2 O 2, acidified with acetic acid:
2KMnO 4 + 5H 2 O 2 + 6CH 3 COOH → 2Mn (CH 3 COO) 2 + 2CH 3 COOK + 8H 2 O + 5O 2
65. Rhenium compounds (II), (III), (VI). Rhenium (VII) compounds: oxide, rhenium acid, perrhenates.
Rhenium(II) oxide- inorganic compound, rhenium metal oxide with the formula ReO, black crystals, insoluble in water, forms hydrates.
Rhenium oxide hydrate ReO H 2 O is formed by the reduction of rhenium acid with cadmium in an acidic medium:
Rhenium(III) oxide- inorganic compound, rhenium metal oxide with the formula Re 2 O 3 , black powder, insoluble in water, forms hydrates.
Obtained by hydrolysis of rhenium(III) chloride in an alkaline medium:
Easily oxidized in water:
Rhenium(VI) oxide- inorganic compound, rhenium metal oxide with the formula ReO 3 , dark red crystals, insoluble in water.
Receipt.
· Proportionation of rhenium(VII) oxide:
Recovery of rhenium(VII) oxide with carbon monoxide:
Chemical properties.
Decomposes on heating:
Oxidized by concentrated nitric acid:
Forms rhenites and perrhenates with alkali metal hydroxides:
Oxidized by atmospheric oxygen:
Recovered with hydrogen:
Rhenium(VII) oxide- inorganic compound, rhenium metal oxide with the formula Re 2 O 7 , light yellow hygroscopic crystals, soluble in cold water, reacts with hot water.
Receipt.
Oxidation of metallic rhenium:
Decomposition on heating of rhenium(IV) oxide:
Rhenium(IV) oxide oxidation:
Decomposition upon heating of rhenium acid:
Chemical properties.
Decomposes on heating:
· Reacts with hot water:
Reacts with alkalis to form perrhenates:
It is an oxidizing agent:
Recovered with hydrogen:
In proportion to rhenium:
Reacts with carbon monoxide:
Rhenic acid- an inorganic compound, an oxygen-containing acid with the formula HReO 4 , exists only in aqueous solutions, forms salts perrhenates.
The transfer of rhenium from poorly soluble compounds, such as ReO and ReS2, into solution is carried out by acid decomposition or alkaline fusion with the formation of soluble perrhenates or rhenium acid. Conversely, the extraction of rhenium from solutions is carried out by its precipitation in the form of slightly soluble perrhenates of potassium, cesium, thallium, etc. Ammonium perrhenate is of great industrial importance, from which metallic rhenium is obtained by reduction with hydrogen.
Rhenic acid is obtained by dissolving Re2O7 in water:
Re2O7 + H2O = 2HReO4.
Solutions of rhenium acid were also obtained by dissolving metallic rhenium in hydrogen peroxide, bromine water, and nitric acid. Excess peroxide is removed by boiling. Rhenic acid is obtained by oxidation of lower oxides and sulfides, from perrhenates using ion exchange and electrodialysis. For convenience, Table 2 shows the density values of rhenium acid solutions.
Rhenic acid is stable. Unlike perchloric and permanganic acids, it has very weak oxidizing properties. Recovery is usually slow. Metal amalgams and chemical agents are used as reducing agents.
Perrhenates are less soluble and thermally more stable than the corresponding perchlorates and permanganates.
Thallium, cesium, rubidium and potassium perrhenates have the lowest solubility.
Perrhenates Tl, Rb, Cs, K, Ag are poorly soluble substances, perrhenates ,Ba, Pb (II) have an average solubility, perrhenates Mg, Ca, Cu, Zn, Cd, etc. dissolve very well in water. In the composition of potassium and ammonium perrhenates, rhenium is isolated from industrial solutions.
Potassium perrhenate KReO4 - small colorless hexagonal crystals. It melts without decomposition at 555°, at higher temperatures it volatilizes, partially dissociating. The solubility of the salt in an aqueous solution of rhenium acid is higher than in water, while in the presence of H2SO4 it remains virtually unchanged.
Ammonium perrhenate NH4ReO4 is obtained by neutralizing rhenium acid with ammonia. Relatively well soluble in water. Upon crystallization from solutions, it forms continuous solid solutions with KReO4. When heated in air, it decomposes starting at 200°C, giving sublimation containing Re2O7 and a black residue of ReO2. When decomposed in an inert atmosphere, only rhenium (IV) oxide is formed according to the reaction:
2NH4ReO4 = 2ReO2 + N2 + 4H2O.
When a salt is reduced with hydrogen, a metal is obtained.
Of the salts of rhenium acid with organic bases, we note nitrone perrhenate C20H17N4ReO4, which has a very low solubility in acetate solutions, especially in the presence of an excess of nitrone acetate. The formation of this salt is used to quantify rhenium.
manganese compounds. Oxides, hydroxides. Permanganic acid. Potassium permanganate, its oxidizing properties in acidic, neutral and alkaline media.
Manganese compounds. Oxides, hydroxides.
Manganese forms several oxides. The most stable are
MnO Mn2O3 MnO2 Mn2O7
Manganese oxide (VII) Mn2O7 is a black-green oily liquid. Above 50 ° C, it decomposes with the formation of oxygen and lower oxides, and explodes at a higher temperature:
2Mn2O7 = 4MnO2 + 3O2.
Shows acidic properties. Reacts with water to form permanganic acid:
Mn2O7 + H2O = 2HMnO4.
Manganese oxide can only be obtained indirectly:
2KMnO4 + H2SO4 = Mn2O7 + K2SO4 + H2O.
Permanganic acid is a strong acid, very unstable, it decomposes already above 3°C:
4HMnO4 = 4MnO2 + 2H2O + 3O2.
Manganese (II) oxide MnO and the corresponding hydroxides Mn (OH) 2 are basic substances.
Οʜᴎ interact with acids to form manganese (II) salts
MnO + 2HCl = MnCl2 + 2H2O
Mn(OH)2 + 2HCl = MnCl2 + 2H2O
Mn(OH)2 is obtained by the action of alkalis on soluble Mn2+ salts
MnCl2 + 2NaOH = Mn(OH)2↓ + 2Н2O
Mn2+ + 2OH- = Mn(OH)2↓
white precipitate
Due to the instability of Mn(OH)2, it oxidizes already in air, forming Mn(OH)4
2Mn(OH)2 + O2 + 2H2O = 2Mn(OH)4
This reaction is qualitative for the Mn2+ cation
Qoxide manganese (IV), or manganese dioxide, MnO2 and hydroxide Mn (OH) 4 are amphoteric substances.
When MnO2 interacts with sulfuric acid, low-stable manganese (IV) sulfate is formed
МnО2 + 2H2SO4 = Mn(SO4)2 + 2 Н2O
When MnO2 is fused with alkalis, the reaction proceeds with the formation of manganites (IV), which should be considered as salts of manganese acid H4MnO4
MnO2 + 4KOH = K4MnO4 + 2H2O
Manganese (IV) oxide, based on the substances with which it reacts, can exhibit the properties of both an oxidizing agent and a reducing agent.
4HCl + MnO2 = MnCl2 + Cl2 + 2H2O
2MnO2 + 3PbO2 + 6HNO3 = 2HMnO4 + 3Pb(NO3)2 + 2 H2O
In the first reaction, MnO2 acts as an oxidizing agent, in the second - as a reducing agent.
Τᴀᴋᴎᴍ ᴏϬᴩᴀᴈᴏᴍ, in the series of oxides and hydroxides of manganese with different degrees of oxidation, a general pattern is manifested: with an increase in the degree of oxidation, the basic character of oxides of hydroxides weakens, and the acid character increases.
Salts of manganese acid are called permanganates.
The most famous is the salt of potassium permanganate KMnO4 - a dark purple crystalline substance, sparingly soluble in water. KMnO4 solutions have a dark crimson color, and at high concentrations - violet, characteristic of MnO4- anions.
Potassium permanganate decomposes when heated
2KMnO4 = K2MnO4 + MnO2 + O2
Potassium permanganate is a very strong oxidizing agent; it easily oxidizes many inorganic and organic substances. The degree of manganese reduction depends very much on the pH of the medium.
Salts of manganese acid - permanganates - contain permanganate ion MnO4-, in solution - purple. They exhibit oxidizing properties, manganese (II) compounds are formed in an acidic environment:
2KMnO4 + 5K2SO3 + 3H2SO4 = 2MnSO4 + 6K2SO4 + 3H2O
in neutral - manganese (IV):
2KMnO4 + 3K2SO3 + H2O = 2MnO2 + 3K2SO4 + 2KOH
in alkaline - manganese (VI):
2KMnO4 + K2SO3 + 2KOH = 2K2MnO4 + K2SO4 + H2O
When heated, they decompose:
2KMnO4 = K2MnO4 + MnO2 + O2.
Potassium permanganate is obtained according to the following scheme:
2MnO2 + 4KOH + O2 = 2K2MnO4 + 2H2O;
then the manganate is converted to permanganate by electrochemical oxidation, the overall process equation is:
2K2MnO4 + 2H2O = 2KMnO4 + 2KOH + H2.
manganese compounds. Oxides, hydroxides. Permanganic acid. Potassium permanganate, its oxidizing properties in acidic, neutral and alkaline media. - concept and types. Classification and features of the category "Compounds of manganese. Oxides, hydroxides. Permanganic acid. Potassium permanganate, its oxidizing properties in acidic, neutral and alkaline media." 2017, 2018.
Author Chemical Encyclopedia b.b. I.L.KnunyantsMANGANESE OXIDES: MnO, Mn 2 O 3, MnO 2, Mn 3 O 4, Mn 2 O 7, Mn 5 O 8. Except Mn 2 O 7, all oxides are crystals, insoluble in water. When higher oxides are heated, O 2 is split off and lower oxides are formed:
When kept in air or in an atmosphere of O 2 above 300 ° C, MnO and Mn 2 O 3 are oxidized to MnO 2.
Anhydrous and hydrated. Mn oxides are part of manganese and iron-manganese ores in the form of minerals pyrolusite b -MnO 2, psilomelan mMO * nMnO 2 * xH 2 O [M \u003d Ba, Ca, K, Mn (H)], manganite b -MnOOH (Mn 2 O 3 * H 2 O), groutite g-MnOOH, brownite 3Mn 2 O 3 * MnSiO 3 and others with MnO 2 content of 60-70%. Processing of manganese ores includes wet enrichment and subsequent chemical separation of oxides MnO 2 or Mn 2 O 3 by sulfitization and sulfatization, carbonization, restore. roasting, etc.
MnO monoxide (mineral manganosite). Up to - 155.3 ° C, hexagon is stable. modification, above - cubic (see table). Semiconductor. Antiferromagnet with Neel point 122 K; magn. susceptibility + 4.85 * 10 - 3 (293 K). Possesses weakly basic properties; is reduced to Mn by hydrogen and active metals when heated. When MnO interacts with acids, Mn(II) salts are formed, with a NaOH melt at 700-800 ° C and an excess of O 2 - Na 3 MnO 4 , with the action of (NH 4) 2 S - MnS sulfide. Obtained by decomposition of Mn (OH) 2, Mn (C 2 O 4), Mn (NO 3) 2 or MnCO 3 in an inert atmosphere at 300 ° C, controlled reduction of MnO 2 or Mn 2 O 3 with hydrogen or CO at 700-900 ° WITH. Component of ferrites and other ceramics. materials, slag for metal desulfurization, microfertilizers, piperidine dehydrogenation catalyst, antiferromagnet. material.
Sesquioxide Mn 2 O 3 exists in two modifications - rhombic. a (mineral kurnakite) and cubic. b (bixbyite mineral), transition temperature a : b 670 °С; paramagnetic, magnetic susceptibility +1.41 10 - 5 (293 K); H 2 is reduced at 300 ° C to MnO, aluminum when heated - to Mn.
Under the action of dilute H 2 SO 4 and HNO 3 passes into MnO 2 and Mn(II) salt. Get Mn 2 About 3 thermodynamic decomposition of MnOOH.
Manganese oxide (II, III) Mn 3 O 4 (mineral hausmanite); a -Mn 3 O 4 at 1160°C goes into b -Mn 3 O 4 with cubic crystalline. lattice; D H 0 transition a : b 20.9 kJ/mol; paramagnetic, magnetic susceptibility + 1.24 * 10 - 5 (298 K). Shows the chemical properties inherent in MnO and Mn 2 O 3 .
MnO 2 dioxide is the most common Mn compound in nature; the most stable b-modification (mineral pyrolusite). Known rhombic. g -MnO 2 (mineral ramsdelite, or polyanite), as well as a, d and e, considered as solid solutions of various forms of MnO 2. Paramagnetic, magnetic susceptibility + 2.28 * 10 - 3 (293 K). Mn dioxide - non-stoichiometric. compound, in its lattice there is always a lack of oxygen. Amphoterene. H 2 is reduced to MnO at 170°C. When interacting with NH 3, H 2 O, N 2 and Mn 2 O 3 are formed. Under the action of O 2 in the melt, NaOH gives Na 2 MnO 4, in conc. acids - the corresponding salts of Mn(IV), H 2 O and O 2 (or Cl 2 in the case of hydrochloric acid). MnO 2 is obtained by decomposition of Mn(NO 3) 2 or Mn(OH) 2 at 200°C in air, reduction of KMnO 4 in a neutral medium, electrolysis of Mn(II) salts. It is used to obtain Mn and its compounds, desiccants, as a depolarizer in dry elements, a component of brown pigment (umber) for paints, for lightening glass, as a reagent for the detection of Cl - , an oxidizing agent in hydrometallurgy Zn, Cu, U, a catalyst component in hopcalite cartridges and other Active MnO 2 , obtained by the interaction of aqueous solutions of MnSO 4 and KMnO 4 , is an oxidizing agent in organic chemistry.
Manganese oxide (VII) Mn 2 O 7 (dimanganese heptaoxide, manganese anhydride) - oily green liquid; melting point 5.9 °C; density 2.40 g/cm 3 ; D H 0 arr -726.3 kJ / mol. Above 50 °C, with slow heating, it begins to decompose with the release of O 2 and the formation of lower oxides, and explodes at higher temperatures or high heating rates; extremely sensitive to mechanical and thermal influences. Strong oxidizing agent; on contact with Mn 2 O 7 combustible substances ignite. MANGANESE OXIDEb. obtained by the interaction of KMnO 4 with H Z SO 4 in the cold.
Oxide Mn 5 O 8, or Mn 2 II (Mn IV O 4) 3, is a solid; insoluble in water; can be obtained by oxidation of MnO or Mn 3 O 4 ; easily decomposes into MnO 2 and O 2 .
From hydroxides Mn stoichiometric. compound are only Mn(OH) 2 , MnO(OH) and HMnO 4 , the others are hydratir. oxides of variable composition, similar in chemical properties to the corresponding oxides. The acidic properties of hydroxides increase with increasing oxidation state of Mn: Mn(OH) 2< MnО(ОН) (или Mn 2 O 3 * xH 2 O) < MnO 2 * xН 2 О < Mn 3 О 4 * xН 2 О < Н 2 MnО 4 < НMnО 4 . Гидроксид Мn(II) практически не растворим в воде (0,0002 г в 100 г при 18 °С); основание средней силы; растворим в растворах солей NH 4 ; на воздухе постепенно буреет в результате окисления до MnО 2 * xН 2 О.
Hydroxyxide Mn(III) MnO(OH) is known in two modifications; at 250 °C in vacuum it is dehydrated to g-Mn 2 O 3 ; in water not sol. Natural Manganite does not decompose HNO 3 and diluted H 2 SO 4 , but slowly reacts with H 2 SO 3 , artificially obtained is easily decomposed by mineral acids; O 2 is oxidized to b-MnO 2. See also Manganates.
MANGANESE OXIDEo. toxic; MPC, see Art. Manganese.
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Aluminothermic method, restoring oxide Mn 2 O 3 formed during pyrolusite calcination: ;
Mn 6+ disproportionates to MnO 2 and MeMnO 4 .
Mn 3+ disproportionate into Mn 2+ and MnO 2 (solid)
Oxides:
Manganese oxides are known, in which it is in the oxidation states +2, +3, +4, +7, as well as several mixed ones, for example Mn 3 0 4 . With an increase in the degree of oxidation, the covalent nature of the Mn-O bond increases, there is a weakening of the basic and strengthening of acidic properties, an increase in oxidative activity. Thus, manganese(2) oxide exhibits the main and predominant reducing properties, and higher manganese oxide has acid
lot character. It has pronounced oxidizing properties. Manganese oxides in intermediate oxidation states are amphoteric and are easily reduced to Mn(II), exhibiting the properties of strong oxidizing agents.
Manganese oxide (2). The lower oxide MnO is a gray-green crystalline powder with the NaCl structure. It occurs naturally as the mineral manganosite. His
the reactivity depends on the grain size. The highly dispersed form of the oxide, formed during low-temperature decomposition (420-450 ° C) of manganese (II) carbonate, oxalate or acetate, has significant chemical activity - it ignites in air. Its calcination in an inert atmosphere leads to coarsening of grains and weakening of chemical activity. To obtain a stoichiometric phase, the Mn 3 0 4 and Mn0 2 oxides are reduced with hydrogen at high temperatures. Lower manganese oxide can also be obtained by the interaction of manganese
with carbon dioxide at a temperature of 300 "C, reduction or calcination of other manganese oxides, as well as during dehydration of Mn (OH) 2 hydroxide in a reducing atmosphere at 800 ° C.
Manganese(2) oxide is predominantly basic in character. It readily dissolves in acids to form divalent manganese salts. Weak amphotericity appears only when
prolonged boiling with concentrated alkali solutions:
MnO + 20H - + H 2 0 \u003d 4 [Mn (OH) 4] 2- Fusion with sodium hydroxide leads to the oxidation of manganese: MnO + 2NaOH \u003d Na 2 Mn0 3 + H 2
Manganese(2) oxide is practically not reduced by hydrogen, but it reacts with active metals. The reaction with sodium proceeds already at room temperature, and aluminothermy requires initiation with an incendiary mixture. Carbon or carbon monoxide can also be used as a reducing agent.
Manganese(2) oxide is a strong reducing agent - at elevated temperatures it reacts even with such a weak oxidizing agent as carbon monoxide(4).
Manganese oxide(3). Oxide Mn 2 0 3 is a brown
powder that decomposes when heated. when stored in air, it slowly oxidizes to manganese (4) oxide, and in the eutectic melt NaOH-KOH (227 ° C), oxygen converts it into
manganate(5). Manganese(3) oxide does not interact with water, but disproportionates in an acidic environment.
Like the lower oxide of manganese, it is predominantly basic in character, but under certain conditions it exhibits some signs of amphotericity. So, with prolonged boiling with alkali solutions, anionic hexahydroxomanganates (3) [Mn (OH) 6 ] 3- are formed, which exist only in a strongly alkaline medium. Mn 2 0 3 exhibits oxidizing properties when boiled with concentrated hydrochloric acid.
Manganese oxide (4). For manganese, this is the most common mineral pyrolusite Mn0 2 found in nature. Among the many Mn0 2 formation reactions, only a few are preferred due to their ease of operation and reproducibility. These include the oxidation of manganese(2) compounds with solutions of chlorates, permanganate
nats, persulfates, chlorine and ozone, reduction of permanganates by the action of hydrogen peroxide, hydrochloric acid, sulfur dioxide, sulfites, alcohols, manganese salts (2). Most of the manganese dioxide is released as a brown precipitate, but part of it remains in the form of a colloidal solution. Manganese dioxide, as a compound of an intermediate oxidation state, exhibits oxidizing and reducing properties. It is more characteristic of the behavior of an oxidizing agent. When heated, an intramolecular redox reaction occurs, accompanied by the release of oxygen and a gradual decrease in the oxidation state of manganese with the successive formation of various oxides. With dry hydrogen sulfide, the reaction proceeds very quickly already at room temperature. When hydrogen sulfide is passed through an aqueous suspension of Mn0 2, a whitish precipitate is formed, consisting of a mixture of manganese (2) sulfide and sulfur; the solution contains S0 4- and S 2 0 2- ions.
The oxidizing properties of manganese dioxide are most pronounced in a strongly acidic environment. When manganese dioxide is treated with cold concentrated hydrochloric acid, a dark solution remains, containing complex acid H 2 [MnC1 6 ], and possibly MnC1 4 tetrachloride not isolated individually. Gradually, the solution brightens due to the reduction of chloride
complexes [MnC1 6 ] 2- -» [MnC1 6 ] 3- -> Mn 2+ + 6C1 - and chlorine evolution.
In nitric and dilute sulfuric acids, Mn0 2 is insoluble, only when it is stored for a long time under an acid solution, a gradual release of oxygen is observed due to the oxidation of water. The interaction of Mn0 2 dioxide with alkalis proceeds only in concentrated solutions and melts. Depending on the reaction temperature, the ratio of reagents, the presence or absence of oxidizing agents, such as atmospheric oxygen, anionic oxomanganates are formed containing manganese in different oxidation states from +3 to +6. For example, when manganese dioxide reacts with a hot concentrated potassium hydroxide solution in an inert atmosphere, a green solution of manganate(4) is first formed, which then turns into a dark blue solution containing equimolar amounts of manganate(5) and manganate(3). When the reaction is carried out in air at a temperature of 350 - 450 ° C, manganate (6) is formed, at a temperature of 600 - 800 ° C - manganate (5).
Hydroxides.
Manganese hydroxide(2), precipitated in the form of a white (ivory) precipitate under the action of alkalis on aqueous solutions of manganese (2) salts in air is rapidly oxidized, so its synthesis is carried out in an inert atmosphere. Stronger oxidizing agents, such as bromine water, hypochlorite, convert it to pyrolusite. Unlike most transition metal hydroxides, hydroxide
manganese(2) has a stoichiometric composition corresponding to the formula Mn(OH)2; it is isostructural with magnesium hydroxide; is a base of medium strength, superior to ammonia, therefore, it goes into solution not only under the action of acids, but also in the presence of ammonium ions.
The weak amphoteric nature of Mn (2) hydroxide is manifested in its ability to form anionic hydroxo complexes: Mn (OH) 2+ OH -<->[Mn (OH) 3] -
Manganese hydroxide(3) is formed in the form of a brown precipitate of variable composition during the oxidation of manganese(2) hydroxide with atmospheric oxygen. The initially formed product contains manganese atoms in two oxidation states: +2 and +3 - and corresponds to the composition MnOx * nH 2 0.
On standing, it transforms into the form of MnOOH oxohydroxide, known in the form of two modifications: orthorhombic (y-MnOOH, manganite) and monoclinic (a-MnOOH, groutite). Manganese(3) oxohydroxide is amphoteric.
Salt.
Various salts are obtained by the interaction of manganese(2) hydroxide with acids. Unlike hydroxide, all of them are resistant to oxidation by atmospheric oxygen.
Salts of manganese(2) form hydrates, painted in a pale pink color, characteristic of [Mn(H 2 0)6] 2+ cations. They are part of some hydrates, for example, MnS0 4 * 7H 2 0, Mn (C10 4) 2 * 6H 2 0, Mn (N0 3) 2 * 6H 2 0. Hydrolysis occurs in aqueous solutions of manganese (2) salts: [Mn (H 2 0) 6] 2+ + H 2 0 \u003d [Mn (H 2 0) 5 OH] + + H 3 0 +
The action of alkali metal sulfites on solutions of manganese salts leads to the precipitation of basic salts, for example NaMn 2 OH (S0 3) 2 (H 2 0).
When Mn 2+ ions are precipitated with sulfide solutions, a solid-colored precipitate of hydrated manganese(2) sulfide MnS *xH20 precipitates. Although the crystal structure of this product is unknown, the pale pink color indicates that water molecules complete the manganese coordination sphere to an octahedron. When stored in air, the precipitate gradually oxidizes, turning into brown oxohydroxide MnOOH, and in an inert atmosphere loses water of crystallization, acquiring a green color characteristic of the natural mineral alabandin (α-MnS), which has the structure
tour of the NaCl type. By acting on a solution of a manganese salt with sodium dihydrogen phosphate in an environment close to neutral, it is possible to precipitate a white precipitate of the average orthophosphate Mn 3 (P0 4) 2 * 7H 2 0.
In an alkaline environment, basic salts precipitate, for example, Mn 2 (OH) P0 4, Mn 5 (OH) 4 (P0 4) 2.
Various acid phosphates are also known, for example, Mn (H 2 P0 4) 2 * H 2 0, highly soluble in water, and slightly soluble MnHP0 4 * H 2 0. When exposed to solutions of manganese (2) salts with a solution of sodium carbonate, a white precipitate of carbonate is released MnC0 3 containing only a small amount of basic salt. To avoid its formation, precipitation is carried out with bicarbonate.
Salts of manganese (3)
When MtOOH is dissolved in the cold in 70% sulfuric acid, a red solution is formed, from which red crystals can be isolated. The most common method for obtaining manganese(3) salts is redox transformations. Thus, manganese(3) sulfate is synthesized by anodic oxidation of a hot solution of manganese(2) sulfate in sulfuric acid or by reduction of potassium permanganate with sulfuric acid:
2KMn0 4 + 4H 2 S0 4 = Mn 2 (S0 4) 3 + K 2 S0 4 + 20 2 + 4H 2 0
Inorganic manganese salts(4) extremely unstable and little studied. Thus, manganese (4) sulfate Mn (S0 4) 2 is isolated in the form of black crystals from sufficiently concentrated sulfuric acid solutions of manganese (2) sulfate when potassium permanganate is added to them.
Salts of manganese (6).
For the first time, manganates (6) were prepared by I. Glauber in
1659 when dissolving pyrolusite in molten saltpeter. These substances are dark green, almost black crystals, soluble in alkali solutions to form emerald green solutions.
In neutral and slightly acidic solutions, manganates are easily dispro-
portioned, as judged by the precipitation of a dark brown precipitate of dioxide and the appearance of a crimson color of the solution. To speed up the process, carbon dioxide is usually passed through the solution. Among the manganates, the most important is the potassium salt K 2 Mn0 4, which is used for the production of KM p 04 permanganate. In laboratory conditions, alkali metal manganates are conveniently obtained
oxidation of water with permanganate in a strongly alkaline solution. To do this, permanganate is placed in a concentrated alkali solution and heated.
Manganese dioxide (MnO 2) or manganese (IV) oxide is a dark gray substance. When heated in air to 530 degrees. C manganese dioxide decomposes releasing oxygen as shown above. In a vacuum or in the presence of a reducing agent, this reaction proceeds much more intensively.
When manganese dioxide is boiled with concentrated nitric acid, a manganese (II) salt is formed and oxygen is released:
2 MnO 2 + 4 HNO 3 \u003d 2 Mn (NO 3) 2 + 2 H 2 O + O 2.
Manganese dioxide in an acidic environment exhibits oxidizing properties:
MnO 2 + 4 Hcl \u003d MnCl 2 + Cl 2 + 2 H 2 O;
MnO 2 + 2 FeSO 4 + 2 H 2 SO 4 \u003d MnSO 4 + Fe 2 (SO 4) 3 + 3 H 2 O.
When manganese (IV) oxide is fused with alkalis without air access, manganite or manganate (IV) is formed:
2 MnO 2 + 2 KOH \u003d K 2 MnO 3 + H 2 O.
In the presence of atmospheric oxygen, which plays the role of an oxidizing agent, a manganate salt (VI) is formed during fusion:
2 MnO 2 + 4 KOH + O 2 \u003d 2 K 2 MnO 4 + 2 H 2 O.
Potassium manganate (K 2 MnO 4) spontaneously decomposes into potassium permanganate and manganese dioxide:
3 K 2 MnO 4 + 2 H 2 O \u003d 2 KMnO 4 + MnO 2 + 4 KOH.
Potassium permanganate (KMnO 4) is widely used in laboratory practice, industry, medicine and everyday life. It is a very strong oxidizing agent. Depending on the environment, manganese in the presence of a reducing agent can be reduced to various degrees of oxidation. In an acidic environment, it is always reduced to Mn (II):
2 KMnO 4 + 10 KBr + 8 H 2 SO 4 \u003d 2 MnSO 4 + 6 K 2 SO 4 + 5 Br 2 + 8 H 2 O.
Potassium manganate (K 2 MnO 4) and manganese dioxide behave similarly.
In an alkaline environment, potassium permanganate is reduced to manganate:
2 KMnO 4 + K 2 SO 3 + 2 KOH \u003d K 2 SO 4 + 2 K 2 MnO 4 + H 2 O.
In a neutral or slightly alkaline environment, potassium permanganate is reduced to manganese dioxide:
2 KMnO 4 + C 6 H 5 CH 3 \u003d 2 KOH + 2 MnO 2 + C 6 H 5 COOH;
2 KMnO 4 + 3 MnSO 4 + 2 H 2 O \u003d 5 MnO 2 + K 2 SO 4 + 2 H 2 SO 4.
The latter reaction is used in analytical chemistry for the quantitative determination of manganese.
Previously, potassium permanganate was obtained by oxidation of either manganese dioxide or potassium manganate. Manganese dioxide was oxidized with saltpeter when fused with alkali:
MnO 2 + KNO 3 + 2 KOH \u003d K 2 MnO 4 + KNO 2 + H 2 O.
The resulting potassium manganate in solution spontaneously decomposed into potassium permanganate and manganese dioxide:
3 K 2 MnO 4 + 2 H 2 O \u003d 2 KMnO 4 + MnO 2 + 4 KOH.
According to the second method, potassium manganate was oxidized with chlorine:
2 K 2 MnO 4 + Cl 2 \u003d 2 KMnO 4 + 2 KCl.
Currently, potassium permanganate is obtained by electrolytic oxidation of manganate:
MnO 4 2- - e - \u003d MnO 4 -.
Potassium permanganate is widely used both in industry and in laboratory practice. It is used to bleach cotton, wool, spinning fibers, clarify oils and oxidize various organic substances. In laboratory practice, it is used to obtain chlorine and oxygen:
2 KMnO 4 + 16 HCl \u003d 2 KCl + 2 MnCl 2 + 5 Cl 2 + 8 H 2 O;
2 KMnO 4 \u003d K 2 MnO 4 + MnO 2 + O 2.
In analytical chemistry, potassium permanganate is used for the quantitative determination of substances with reducing properties (Fe 2+, Sn 2+, AsO 3 3+, H 2 O 2, etc.). This method of analysis is called permanganatometry.
