Chambers, Holliday. Modern inorganic chemistry

376 THE TRANSITION ELEMENTS

one electron required (and also giving the vanadium a formal oxidation state of — 1).

TESTS FOR VANADIUM

The colour sequence already described, for the reduction of vanadium(V) to vanadium(II) by zinc and acid, gives a very characteristic test for vanadium. Addition of a few drops of hydrogen peroxide to a vanadate(V) gives a red colour (formation of a peroxo-complex) (cf. titanium, which gives an orange-yellow colour).

CHROMIUM

THE ELEMENT

Chromium occurs quite extensively, mainly as the ore chromite or chrome ironstone, a mixed oxide of iron(II) and chromium(III). Presence of chromium in the mineral beryl produces the green colour of emeralds and the red colour of ruby is due to the substitution of Cr(III) for Al(III)in the mineral alurmnium(III) oxide; hence the name ^chromium' derived from the Greek for colour. Direct reduction of chromate by heating with carbon and calcium oxide gives an alloy of iron and chromium, ferrochrome, which can be added to steel, to make stainless steel (12-15 % chromium). The pure metal can be prepared by reduction of the + 3 oxide, Cr2O3, using powdered aluminium, or by electrolytic reduction of the + 6 oxide CrO3. The metal is extensively used in chromium plating because it is relatively inert to chemical attack. However, the extent of inertness is dependent on its purity. It is inert to the oxidising oxo-acids (phosphoric, nitric, aqua-regia, concentrated sulphuric); these render it passive, probably by formation of a surface layer of oxide. It remains bright in air, despite formation of a surface layer of oxide. When pure (no oxide layer) it is readily soluble in dilute hydrochloric acid (to give a chromium(II) cation, see below) and displaces copper, tin and metal from solutions of their salts.

In the older form of the periodic table, chromium was placed in Group VI, and there are some similarities to the chemistry of this group (Chapter 10). The outer electron configuration, 3d5 4s1. indicates the stability of the half-filled d level, 3d5 4sl being more stable than the expected 3d4 4s2 for the free atom. Like vanadium and titanium, chromium can lose all its outer electrons, giving chromium(VI); however, the latter is strongly oxidising and is

THE TRANSITION ELEMENTS 377

therefore only found in combination with oxygen and fluorine. Of the lower oxidation states, the + 3 is the most stable and common.

COMPOUNDS OF CHROMIUM

Oxidation state + 6

In this state, chromium compounds are usually coloured yellow or red (but due to charge transfer (p. 60) and not to the presence of d electrons on the chromium ion). The only halide known is the unstable chromium(VI) fluoride CrF6, a yellow solid. However, oxide halides are known, for example CrO2Cl2 Cchromyl chloride'), formed as a red vapour when concentrated sulphuric acid is added to a chromate(VI) (or dichromate) mixed with a chloride:

Cr2O?- -h 4Cr + 6H2SO4 -> 2CrO2Cl2 + 6HSO4 + 3H2O

(This reaction may be used to distinguish a chloride from a bromide, since CrO2Br2 is unstable under these conditions).

The most important compounds containingCr(VI) are the oxide CrO3 and the oxoanions CrOj", chromate(VI) and C^O^". dichromate(VI).

CHROMIUM(VI) OXIDE (CHROMIUM TRIOXIDE)

Chromium trioxide is obtained as bright red crystals when concentrated sulphuricacid is added cautiously to a concentrated aqueous solution of a chromate or dichromate(VI). It can be filtered off through sintered glass or asbestos, but is a very strong oxidising agent and so oxidises paper and other organic matter (hence the use of a solution of the oxide—"chromic acid'—as a cleansing agent for glassware).

Chromium(VI) oxide is very soluble in water ; initially, "chromic acid', H2CrO4, may be formed, but this has not been isolated. If it dissociates thus:

H2 CrO4 ^H+ + HCrO;

then the HCrO^ ions probably form dichromate ions :

^ C r O ~ + HO

Chromium(VI) oxide is acidic, and the corresponding salts are the chromates and dichromates, containingthe ions CrO|~ and Cr2O7 ", i.e. [CrO4 4- CrO3]2~. The oxidation state of chromium is +6 in each ion (cf.sulphur in SO^" and S2 O7~).

378 THE TRANSITION ELEMENTS

THECHROMATES(Vl)

The chromates of the alkali metals and of magnesium and calcium are soluble in water; the other chromates are insoluble. The chromate ion is yellow, but some insoluble chromates are red (for example silver chromate, Ag2CrO4). Chromates are often isomorphous with sulphates, which suggests that the chromate ion, CrOj", has a tetrahedral structure similar to that of the sulphate ion, SO|~ Chromates may be prepared by oxidising chromium(III) salts; the oxidation can be carried out by fusion with sodium peroxide, or by adding sodium peroxide to a solution of the chromium(III) salt. The use of sodium peroxide ensures an alkaline solution; otherwise, under acid conditions, the chromate ion is converted into the orangecoloured dichromate ion:

and certain metal ions

The dichromate ion has the following geometrical structure (single lines not necessary implying single bonds):

i.e. two tetrahedral CrO4 groups joined by a common oxygenatom. If a metal ion of an insoluble chromate is added to a solution containing the dichromate ion, the chromate is precipitated; for

example with a soluble lead(II) salt:

2Pb2 + + Cr2O?- + H2O -» 2PbCrO4l + 2H+

yellow precipitate of leadchromate

Sodium dichromate is prepared on the large scale by heating powdered chromite with sodium carbonate, with free access of air; the sodium chromate first formed is treated with acid:

4FeCr2O4 4- 8Na2CO3 4- 7O2 -> 8Na2CrO4 + 2Fe2O3 + 8CO2T

2Na2CrO4 + H2SO4 -> Na2SO4 + Na2Cr2O7 + H2O

Sodium sulphate crystallises out in hydrated form (common ion effect) and is filtered off; on concentration, sodium dichromate is obtained. For analytical purposes, the potassium salt. K2Cr2O~. is preferred; potassium chloride is added and the less soluble potassium dichromate obtained.

The dichromate ion is a useful oxidising agent in acid solution, and is used in volumetricanalysis:

380 THE TRANSITION ELEMENTS

Addition of hydrogen peroxide to a solution of a dichromate yields the blue colour of 'peroxochromic acid\ This is a test for soluble chromates and dichromates.

Chromates and dichromates are used in industry as oxidising agents, for example in the coal tar industry, in the leather industry (chrome tanning), and in the dye industry as mordants. Some chromates are used as pigments, for example those of zinc and lead, Chromates and dichromates are poisonous.

Oxidation state + 3

This is the most common and stable state of chromium in aqueous solution. The Cr3 + ion, with 3d3 electrons, forms mainly octahedral complexes [CrX6], which are usually coloured, and are kinetically inert, i.e. the rate of substitution of X by another Ugand is very slow; consequently a large number of such complexes have been isolated (see below, under chromium(III) chloride).

CHROMIUM(IH) CHLORIDE, CrCl3

Chromium(III) chloride is prepared in the anhydrous form:

1. By the reaction of chlorine with a heated mixture of chrom ium(III) oxide and carbon:

Cr2O3 + 3C12 + 3C -> 3COT + 2CrCl3

2. By the reaction of sulphur dichloride oxide with the hydrated chloride:

CrCl3.6H2O 4- 6SOC12 -» CrCl3 + 6SO2t + 12HC1T

Anhydrous chromium(III) chloride is a peach-coloured solid, which is insoluble in water unless a trace of reducing agent is present. Solution then occurs readily to give a green solution from which the green hydrated chloride, CrCl3,6H2O, can be crystallised out. If this substance is treated with silver nitrate, only one third of the chlorine is precipitated; hence the formula is [Crm(H2O)4Cl2]+Cr .2H2O. with two chloride ions as ligands in the complex ion. Two other forms of formula CrCl3.6H2O are known; one is (pale green) LCr(H2O)5Cl]2 + [Cr]2 .H2Olrom which silver nitrate precipitates two thirds of the chlorine; and the other is [Cr(H2O)6]Cl3 (grey-blue) from which all the chlorine is precipitated by silver nitrate. These three compounds are isomers, and the cations can be represented thus:

382 THE TRANSITION ELEMENTS

Chromium(III) hydroxide, like aluminium hydroxide, possesses (Wsorptive power, and the use of chromium compounds as mordants is due to this property.

Chromium(III) hydroxide dissolves in acids to form hydrated chromium(III) salts; in concentrated alkali, hydroxo-complexes [Cr(OH)6]3~ are formed.

CHROMIUM(III) SULPHATE. Cr2(SO4)3.18H.O

Hydrated chromium(III) sulphate exhibits different colours and different forms from which varying amounts of sulphate ion can be precipitated by barium chloride, due to the formation of sulphatocomplexes. Chromium(III) sulphate can form alums.

HYDRATED CHROMIUM(lIl) NITRATE, Cr(NO3)3 .9H2O

Hydrated chromium)III) nitrate is a dark green, very deliquescent solid, very soluble in water. The anhydrous nitrate is covalent.

Oxidation state + 2

This state is strongly reducing, often coloured, and paramagnetic.

384 THE TRANSITION ELEMENTS

TESTS FOR CHROMATES AND DICHROMATES(VI)

1.Addition of lead(II) nitrate in ethanoic acid solution gives a yellow precipitate of lead chromate. PbCrO4.

2.A reducing agent (for example sulphur dioxide) reduces the yellow chromate or orange dichromate to the green chromium(III) state.

3.Hydrogen peroxide with a chromate or a dichromate gives a

blue colour.

TESTS FOR CHROMIUM(III) SALTS (i.e. FOR HYDRATED

Cr3+ ION) :

1.Addition of alkali gives a green gelatinous precipitate of chromium(III) hydroxide, soluble in a large excess of strong alkali.

2.Addition" of sodium peroxide to a solutiongives a yellow colour of the chromate.

MANGANESE

THE ELEMENT

Manganese is the third most abundant transition metal, and is widely distributed in the earth's crust. The most important ore is pyrolusite, manganese(IV) oxide. Reduction of this ore by heating with aluminium gives an explosive reaction, and the oxide Mn3O4 must be used to obtain the metal. The latter is purified by distillation in vacuo just above its melting point (1517 K); the pure metal can also be obtained by electrolysis of aqueous manganese(II) sulphate.

The metal looks like iron; it exists in four allotropic modifications, stable over various temperature ranges. Although not easily attacked by air, it is slowly attacked by water and dissolves readily in dilute acids to give manganese(II) salts. The stable form of the metal at ordinary temperatures is hard and brittle—hence man ganese is only of value in alloys, for example in steels (ferroalloys) and with aluminium, copper and nickel.

COMPOUNDS OF MANGANESE

Although it exhibits a wide range of oxidation states, from -f 1 (corresponding to formal loss of all the outer electrons, 3d54s2) to 0.

THE TRANSITION ELEMENTS 385

it differs from the preceding transition metals in having a very stable 4- 2 oxidation state, corresponding to loss of only the 4s2 electrons, and indicative of the stability of the half-filled d levels.

Oxidation state + 7

Apart from two unstable oxide halides, MnO3F and MnO3Cl, this state is exclusively represented by the oxide Mn2O7 and the anion

MANGANESE(VII) OXIDE, DIMANGANESE HEPTOXIDE, Mn2O7

This oxide is obtained by adding potassium manganate(VII) to concentrated sulphuric acid, when it appears as a dark coloured oil which readily decomposes (explosively on heating) to manganese(IV) oxide and oxygen:

2KMnO4 + 2H2SO4 -> Mn2O7 + 2KHSO4 + H2O 2Mn2O7 -» 4MnO2 + 3O2

It is a powerful and violent oxidising agent. It dissolves in water, and manganic(VII) acid (permanganic acid) HMnO4 and its dihydrate HMnO4.2H2O can be isolated as purple solids by low temperature evaporation of the frozen solution. Manganic(YII) acid is also a violent oxidising agent, especially with any organic material; it decomposes quickly at 276 K.

THE MANGANATES(VII)

The purple manganate(VII) or permanganate anion, MnO^ is tetrahedral; it owes its intense colour to charge transfer (since the manganese has no d electrons). The potassium salt KMnO4 is the usual form, but many other cations from soluble manganate(VII) salts (all purple); those with large unipositive cations (forexample Cs"^) are less soluble. Potassium manganate(VII) can be prepared by (a) electrolytic oxidation of manganese metal (oxidation from 0 to +7) using a manganese anode in potassium carbonate solution,

(b) oxidation of manganate(II) (oxidation + 2 to + 7), using the peroxodisulphate ion S2 Og~* and a manganese(II) salt, and (c)

* This ion oxidises thus:

S2O^ ~ + 2e~ -+ 2SOi ~ : £^ = 2.0 V