Chambers, Holliday. Modern inorganic chemistry

386 THE TRANSITIONELEMENTS

oxidation of manganese(IV), by fusion of MnO2 with potassium hydroxide, the usual method. This fusion, in air or in the presence of a solid oxidising agent (KNO3) produces manganate(VI) ( + 4 to + 6):

2MnO2 + 4KOH + O2 -» 2K2MnO4 + 2H2O

The green manganate(YI) is extracted with water, then oxidised to manganate(VII). This is usually carried out electrolytically, at an anode, but in the laboratory chlorine may be used :

C1 -* 2MnO + 2C1

(Note that here "chlorine' is oxidising the manganate(VI) to manganate(VII) ; under more acid conditions, the latter oxidises chloride to chlorine, p, 103).

Potassium manganate(VII) disproportionates on heating :

2KMnO4 -» K2MnO4 + MnO2 + O2

+ 7 +6 +4

The manganate(VII) ion slowly oxidises water, the essentialreaction being

4MnO4 + 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

MnO^aq) + 8H3O+ + 5e~ -» Mn2+(aq) + 12H2O : E^= + 1.52V

Hence manganate(VII) is used in acid solution to oxidise, for example,

Fe(III), NO2" -* NOj, H2O2 ^ O2,C2Ol" -. 2CO2

quantitatively; the equivalence point is recognised by persistence of the purple colour. (Sulphuric acid is used to acidify, since hydrochloric acid is oxidised to chlorine, and nitric acid is an oxidising agent.) Manganate(VII) is also used extensively in organic chemistry. lor example to oxidise alcohols to aldehydes ; here it may be used in acid or (morecommonly) in alkaline solution, when manganesedV) oxide is the product :

(aq) 4- 2H2O + 3^~ -> MnO2(s) + 4OH~ (aq) : E^ = -h 0.59 V

THE TRANSITION ELEMENTS 387

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

4MnOj + 4OH~ -> 4MnO4~ + 2H2O -f O2

(cf. the reverse reaction with chlorine, above).

Oxidation state + 6

This is only found in the green manganate(VI) ion, already described. It is only stable in alkaline conditions; in neutral or acid solution it disproportionates:

3MnO|~ + 2H2O -* MnO2 + 2MnO4 + 4OH~

Oxidation state + 5

This state exists as a manganate(V), the blue MnO^~ ; Na3MnO4. 10H2O is isomorphous with Na3VO4 (p. 374).

Oxidation state + 4

MANGANESE(IV) OXIDE, MnO2

Manganese(IV) oxide is the only familiar example of this oxidation state. It occurs naturally as pyrolusite, but can be prepared in an anhydrous form by strong heating of manganese(II) nitrate:

Mnn(NO3)2 -> MnO2 4- 2NO2t

It can also be precipitated in a hydrated form by the oxidation of a manganese(II) salt, by, for example, a peroxodisulphate:

Mn2+ + S2Ol" + 2H2O -> 2SO|~ + MnO2l + 4H+

Manganese(IV) oxide is a dark-brown solid, insoluble in water and dilute acids. Its catalytic decomposition of potassium chlorate^) and hydrogen peroxide has already been mentioned. It dissolves slowly in alkalis to form manganates(IV), but the constitution of these is uncertain. It dissolves in ice-cold concentrated hydrochloric acid forming the complex octahedral hexachloromanganate(IV) ion:

MnO2 + 6HC1 -> [MnlvCl6]2~ + 2H+ + 2H2O

388 THE TRANSITION ELEMENTS

This ion is derived from manganese(IV) chloride, MnCl4, but the latter has not been isolated. The MnCl^" ion is unstable,breaking down to give chlorine thus:

[MnClJ2"^ Mn2+ + 4CP + C12T

Hence, under ordinary conditions, manganese(IV) oxide oxidises concentrated hydrochloric acid to chlorine, but the above shows that the oxidation process is essentially :

Mnlv + 2e~ -> Mn°

An oxidation which can be used to estimate the amount of manganese(IV) oxide in a sample of pyrolusiteis that of ethanedioic acid :

MnO2 + (COOH)2 + H2SO4-> MnSO4 + 2CO2T + 2Hf O

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 thiosulphate:

MnO2=Cl2=I2=2S2Or

Manganese(IV) oxide is used as a depolariser in Leclanche cells (the cells used in ordinary batteries), as a glaze for pottery and as a decoloriser for glass. The decolorising action occurs because the manganese(IV) oxide oxidises green iron(II) silicates to the less evident iron(III) compounds; hence the one-time name of "glassmaker's soap' and also "pyrolusite' (Greek pur and lusis, dissolution by fire).

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 + 3

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

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

However the Mn3+(aq) ion can be stabilised by using acidsolutions or by complex formation ;it can be prepared by electrolyticoxidation of manganese(II) solutions. Thealum CaMn(SO4)2 . 12H2O contains

THE TRANSITION ELEMENTS 389

the hydrated Mn3+ ion, which (as expected for a tripositive cation), 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+ 4- 4H2O

Mn3+ + H2O + 5F" -» [MnF5(H2O)]2~

Oxidation of manganese(II) hydroxide by air gives the brown hydrated oxide Mn2O3.aq, and this on drying gives MnO(OH) which occurs in nature as manganite. (The oxide Mn2O3 alsooccurs 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.

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 in a complex of manganese(II) must involve the pairing of electrons, a process which requires energy. Hence electron transitions between the split d-orbitals are weak for manganese(II), and the colour is correspondingly pale (usuallypink). The stability of the d5 configuration with respect to either loss or gain of electrons also means that manganese(II) salts are not easily reduced or oxidised. Indeed, in oxidation state 4-2, manganese shows fewer 'transition-like' characteristics than any other transition metal ion; thus the aquo-ion [Mn(H2O)6]2+ is barely acidic, allowing formation of 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.xH2O and double salts, for example (NH4)2Mn(SO4)2.6H2O; dehydration of the simple hydrated salts, by heating, produces the anhydrous salt withoutdecomposition. Addition of alkali precipitates the white basic manganese(II) hydroxide Mn(OH)2; if left in the alkaline medium it is oxidised readily by air to brown Mn3O3.aq*.

* In water pollution studies, the oxygen content can be measured by making the water alkaline and shaking a measured volume with an oxygen-free solution containing Mn2 ^(aq). The solution is acidified with sulphuric acid, potassium iodide added and the liberated iodine titrated with sodium thiosulphate.

390 THE TRANSITION ELEMENTS

The oxide MnO is obtained by heating the carbonate MnCO3. Oxidation of manganese(II) in aqueous acid solution requires a strong oxidising agent, for example

MnO;(aq) + 8H3O+ + 5e' -* Mn2+(aq) + 12H2O : £e = 1.52 V MnO2(s) + 4H3O+ +2e~ ^Mn2+ (aq)-h6H2 O: £^+ 1.35 V

However, the Mn(II) ion forms a variety of complexes in solution, some of which may be more easily oxidised ; these complexes can be either tetrahedral, for example [MnCl4]2", or octahedral, for example [Mn(CN)6]4'. 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 + .

Low oxidation states

Manganese forms a decacarbonyl Mn2(CO)10 in which each manganese has the required share in 18electrons 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)^ the noble gas structure is again attained.

TESTS FOR MANGANESE

Fusion of a manganese compound with sodium carbonate and potassium nitrate (on porcelain) givesa green manganate(YI) (p. 386).

TEST FOR MnO4 ION

The purple colour of this ion alone is a sufficient test for its presence: addition of sulphuric acid and hydrogen peroxide discharges the colour.

TEST FOR Mn2 + IONS

If a manganese(II) salt is boiled with a strong oxidising agent such

THE TRANSITION ELEMENTS 391

as a peroxodisulphate or lead*IV) oxide and concentrated nitric acid, the purple colour of the manganate(VII) ion is seen.

IRON

THE ELEMENT

After aluminium, iron is the most abundant metal; and the fourth most abundant of all the elements; it occurs chiefly as oxides (for example haematite (Fe2O3), magnetite (lodestone) (Fe3O4) and as iron pyrites FeS2. Free iron is found in meteorites, and it is probable that primitive man used this source of iron for tools and weapons. The extraction of iron began several thousand years ago, and it is still the most important metal in everyday life because of its abundance and cheapness, and its ability to be cast, drawn and forged for a variety of uses.

The process of extraction requires first smelting (to obtain the crude metal)and then refining. In smelting, iron ore (usually an oxide) is mixed with coke and limestone and heated, and hot air (often enriched with oxygen) is blown in from beneath (in a blast furnace). At the lower, hotter part of the furnace, carbon monoxide is produced and this is the essential reducing agent The reduction reactions occurring may be represented for simplicity as:

700 K; by Le Chateliers Principle, more iron is produced higher up the furnace (cooler) than below (hotter). In the hotter region (around.900 K), reaction (13.5) occurs irreversibly, and the iron(II) oxide formed is reduced by the coke [reaction (13.6)] further down. The limestone forms calcium oxide which fuses with earthy material in the ore to give a slag of calcium silicate; this floats on the molten iron (which falls to the bottom of the furnace) and can be run off at intervals. The iron is run off and solidified as "pigs'—boat-shaped pieces about 40 cm long.

Pig-iron or cast iron contains impurities, chiefly carbon (up to 5 %). free or combined as iron carbides. These impurities, some of which form interstitial compounds (p. 1 1 3 ) with the iron, make it hard and brittle, and it melts fairly sharply at temperatures between 1400 and 1500 K; pure iron becomes soft before it melts (at 1812 K). Hence cast iron cannot be forged or welded.

392 THE TRANSITION ELEMENTS

When iron is refined, the process is essentially one of melting the iron in presence of materials which will react with the impurities— for example air (or oxygen) to remove chiefly carbon, and calcium oxide (added as carbonate) to remove phosphorus. There are a variety of refining processes, each depending on the composition of the initial iron and the sort ofiron or steel destined as the end product. Steels have a carbon content of 0.1-1.5%, and addition of other transition metals imparts certain properties (for example a little manganese, elasticity and high tensile strength; more manganese, great hardness; chromium, resistance to chemical attack, as in stainless steel; nickel, a reduced expansion; tungsten and vanadium, hardness retained at high temperatures).

Pure iron is prepared by reduction of iron(II) oxide with hydrogen, or by electrolysis of an iron(II)-containing aqueous solution. It is a fairly soft metal, existingin different form according to temperature:

(It should be noted that the magnetic properties of iron are dependent on purity of the iron and the nature of any impurities.)

Iron combines with most non-metals on heating, and forms the oxides Fe2O3 and (mainly) Fe3O4 when heated in air above 430 K. Steam above 800 K produces the oxide Fe3O4 and hydrogen. Iron dissolves in most dilute acids, giving iron(II) solutions, i.e.

Fe + 2H+ (aq) -> Fe2 + (aq) + H2

This follows from the E^ value for the half-reaction Fe2+(aq) -f 2e" -> Fe(s): £e = - 0.44V

(The impurities in ordinary iron assist dissolution in acid, and are responsible for the characteristic smell of the hydrogen from this source.) In dilute nitric acid, ammonium nitrate is formed:

4Fe + 10H+ 4- NO3 -* 4Fe2+ + NH^ + 3H2O

Concentrated nitric acid renders the metal "passive; i.e.chemically unreactive, due to formation of a thin oxide surface film (which can be removed by scratching or heating inhydrogen).

Iron is a good reducing agent (see the £° value just given): it

THE TRANSITION ELEMENTS 393

reduces some cations to the metal (for example copper) in aqueous solution, giving iron(II).

Iron absorbs hydrogen readily and is a hydrogenation catalyst. In Mendeleef s form of the periodic table, iron (together with cobalt and nickel) was placed in Group VIII and the three elements together were called fca transitional triad'. Hence there was no resemblance to any of the elementsin the main Groups I-VII; these triad elements have properties which are similar, and which show some resemblances to the earlier transition metal properties. However, unlike manganese and the preceding transition elements, iron does not show the maximum possible oxidation state +8 corresponding to the removal of all its eight outer electrons (3d64s2}: the actual maximum oxidation state is +6, but oxidation states above -1-3 are not very important, and +3 and + 2 are the pre dominant and important states for iron. (Cobalt and nickel simi-

larly do not show high oxidation states.)

Oxidation states above + 3

As might be expected, these higher oxidation states are found almost exclusively in anionic form, and are produced only under strongly oxidising conditions.

Alkali metal ferrates(VI), for example K2FeO4, are obtained by oxidation of a suspension of hydrous iron(III) oxide (assumed to be Fe(OH)3 in the equation below)by chlorate(I)in concentrated alkali:

2Fe(OH)3 4- 3C1CT + 4OH~ -* 2FeOr + 3C1" + 5H2O

The deep red FeOj" is stable only in alkali; in acid, iron(III) is produced :

2FeO*- + 1OH+ ^2Fe3+ (aq) + 5H2O + |O2

Ferrate(VI) has powerful oxidising properties, for example ammonia is oxidised to nitrogen. Potassium ferrate(VI) is isomorphous with potassium chromatefVT), and both anions are tetrahedral.

Decomposition of potassium ferrate(VI) at 1000K gives a ferrate(V), K3FeO4, and several types of ferrate(IV), for example FeO|~, FeOt' are known; these ferrates(IV) have no solution chemistry and are probably best regarded as mixed oxides, since the FeOl" ion has no identifiablestructure.

Oxidation state + 3

In this state, iron has five d electrons, but does not show any strong

394 THE TRANSITION ELEMENTS

resemblance to manganese(II), except that most iron(III) compounds show high paramagnetism, i.e. the electrons remain unpaired.

Iron(III) chloride is a black, essentially covalent solid, in which each iron atom is surrounded octahedrally by six chlorine atoms. It is prepared by direct combination of iron with chlorine or bydehydration of the hydrated chloride, by one of the methods given on p. 343).

When the anhydrous solid is heated, it vaporises to form first Fe2Cl6 molecules, then the monomer FeCl3 and finally FeCl2 and chlorine. It fumes in air (with hydrolysis) and dissolves readily in water to give a yellow (dilute) or brown (concentrated) solution, which is strongly acidic. Crystallisation gives the yellow hydrate FeCl3.6H2O which has the structure [FeCl2(H2O)4]CL2H2O, i.e. contains the octahedral complex ion [FeCl2(H2O)4]+ ; ions of this general type are responsible for the colours of the aqueous solution of iron(III) chloride. In the presence of excess chloride *lon, both tetrahedral [FeQ4]~ and octahedral [FeCl6]3~ can be formed.

Iron(III) chloride forms numerous addition compounds, especially with organic molecules which contain donor atoms, for example ethers, alcohols, aldehydes, ketones and amines. Anhydrous iron(III) chloride is soluble in, for example, ether, and can be extracted into this solvent from water; the extraction is more effective in presence of chloride ion. Of other iron(III) halides, iron(III) bromide and iron(III) iodide decompose rather readily into the +2 halide and halogen.

IRON(III) OXIDES AND HYDROXIDE

If an aqueous solution of an iron(III) salt is treated with alkali, a red-brown precipitate of Iron(III) hydroxide' is obtained; this is probably best represented as FeO(OH). On strong heating it gives the red oxide Fe2O3. Iron(III) oxide, Fe2O3, occurs naturally as haematite, and can also be prepared by strong heating of iron(II) sulphate:

2FeSO4 -» Fe2O3 + SO2 + SO.,

It shows some amphoteric behaviour, since it dissolves in alkali (concentrated aqueous or fused) to give a ferrate(III) ; the equation may be written as

FeO + 2OH"

Iron(II) oxide exists in two forms, the red a-form (paramagnetic) and the y-form (ferromagnetic) obtained by careful heating of

THE TRANSITION ELEMENTS 395

kFeO(OH)'. The a-form is used as a red pigment, as a metal polish ("jeweller's rouge') and as a catalyst.

The mixed oxide Fe3O4 (tri-iron tetroxide) is a black solid, which occurs naturally as magnetite; it is formed when iron(III) oxide is

Fe(ClO4)3.lOH2O. In many other salts, the anion may form a complex with the iron(III) and produce a consequent colour change. for example iron(III) chloride hydrate or solution, p. 394. Stable anionic complexes are formed with a number of ions, for example with ethanedioate (oxalate), C2O4~, and cyanide. The redox potential of the ironll-ironlll system is altered by complex formation with each of these ligands; indeed, the hexacyanoferrate(III) ion, [Fe(CN)6]3", is most readily obtained by oxidation of the corresponding iron(II) complex, because

[Fe(H2O)6]3+ + e" -^ [Fe(H2O)6]2+ : E^ = + 0.77 V

[Fe(CN)6]3- + e' -^ [Fe(CN)6]4~ :E^ = + 0.36 V

The thiocyanate ion SCN~ forms an intensely red-coloured complex (most simply represented as [Fe(SCN)(H2O)5]2+) which is a test for iron(III). However, unlike eobalt(III), iron(III) does not form stable hexammines in aqueous solution, although salts containing the ion [Fe(NH3)6]3+ can be obtained by dissolving anhydrous iron(III) salts in liquid ammonia.

Oxidation state -f- 2

In this oxidation state, iron is quite readily oxidised by mild oxidising agents, and hence in many of the reactions it is a mild reducing agent. For acid conditions

Fe3+ (aq) + e~ -> Fe2+(aq): E^ = + 0.77 V

and hence air (oxygen)will be expected to oxidise the + 2 to the +3 state. In practice, this process is usually slow, but more powerful oxidising agents (e.g. manganate(VII) ion, dichromatefVl) ion, hydrogen peroxide) act more rapidly and quantitatively. However this applies strictly only to the green hexaquo-ion [Fe(H2O)6]2* ; a change to higher pH, i.e. to more alkaline conditions,changes the