Chambers, Holliday. Modern inorganic chemistry

396 THE TRANSITION ELEMENTS

+ 2 species finally to insoluble kFe(OH)2' (or hydrated oxide): for this

Fe(OH)3 + e~ -+ Fe(OH)2 + OH'(aq): £^ = - 0.56V

and hence the reducing power is greatly increased, and Te(OH)2* (white when pure) is rapidly oxidised by air. Again, replacement of the water ligands of [Fe(H2O)6]2 + by other ligands will alter the value of E^ (see below, p. 397).

THE HALIDES

The anhydrous halides FeX2 are pale-coloured solids (FeCl2 is yellow) with very high melting points. The chloride may be obtained by heating the metal in a stream of dry hydrogen chloride; it shows some solubility in organic liquids and may be a partly cotalent solid. However, all the halides are deliquescent, and very readily form hydrates. Thus iron(II) chloride forms FeCl2.4H2O and FeCl2,6H2O (both green); in the latter, there are neutral complexes [FeCl2(H20)4].

THE OXIDES

Iron(II) oxide FeO is prepared by heating iron(II) ethanedioate (oxalate) in vacua:

FeC2O4 -» FeO + CO + CO2

It is a black powder, often pyrophoric, and is non-stoichiometric, the formula Fe0 95O more correctly representing its average composition.

The 'hydroxide, Fe(OH)2' has been referred to above.

OTHER IMPORTANTCOMPOUNDS

Other iron(II) salts include, notably the green sulphate heptahydrate FeSO4. 7H2O which on heating yields first the white anhydrous salt FeSO4 and then decomposes :

2FeSO4 -* Fe2O3 + SO2 + SO3

Double salts of general formula M^SO4.FeSO4.6H2O (M = alkali metal or ammonium) can be obtained by crystallisation of solutions containing the appropriate proportions of the two simple salts:

THE TRANSITION ELEMENTS 397

an acid solution of the salt with M = NH4 (Mohr's salt, terrous ammonium sulphate') is much less quickly oxidised by air than is the simple iron(II)sulphate solution, and hence is used in volumetric analysis. Iron(II)sulphide, FeS, may be prepared by heating the elements together, or by precipitation from an iron(II) solution by sulphide ion; it is a black solid which is non-stoichiometric, like the oxide. The yellow sulphide FeS2 (made up essentially of Fe2 + and 82 ~ ions) occurs naturally as pyrites.

COMPLEXES

As with the + 3 state, iron(II) forms a variety of complexes which are usually 6-coordinate and octahedral. Replacement of the water ligands in green [Fe(H2O)6]2+ (itself an octahedral complex) by ammonia molecules is incomplete in aqueous ammonia, but reaction of the anhydrous chloride with gaseous or liquid ammonia gives the complex [Fe(NH3)6]Cl2. The water ligands are more easily replaced by cyanide ions to give the hexacyanoferrate(II) ion, [Fe(CN)6]4~. Many salts of this ion are known, for example the soluble yellow hydrate K4[Fe(CN)6].3H2O, and the insoluble brown copper(II) salt Cu2[Fe(CN)6] once much used as a semi permeable membrane in osmotic pressure determinations. The reaction between aqueous Fe3+ ions and [Fe(CN)6]4~ yields an intense blue precipitate, prussian blue, which is iron(III) hexacyanoferrate(II), Fe4[Fe(CN)6]3; the same material, called TurnhuWs blue. is obtained by addition of Fe2+ (aq.)ions to [Fe(CN)6]3" ions. The intense colour of this compound is due to charge-transfer (p. 60). The formation of [Fe(CN)6]4~ ions causes the iron(II) to change its properties (for example it is not precipitated as the hydroxide with alkali or as the sulphide with S2 ~ ions); it is more readily oxidised to the + 3 state, since

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

When concentrated sulphuric acid is added to a nitrate in the presence of aqueous iron(II) sulphate, the nitrogen oxide liberated forms a brown complex [Fe(H2O)5NO]2+ which appears as a "brown ring' at the acid-aqueous interface (test for a nitrate, p 243).

Perhaps the most important complex of iron(II) is heme (or haeme). Haemoglobin, the iron-containing constituent of the blood, consists essentially of a protein, globin, attached through a nitrogen atom at one coordination position of an octahedral complex of iron(II). Of the other five coordination positions, four (in a plane) are occupied by nitrogen atoms, each of which is part of an organic

398 THE TRANSITION ELEMENTS

OH2

~^v

02

rim: system—the whole system is a porphin. The sixth position (Figure 13.4} is occupied either by an oxygen molecule or a water molecule, and here reversible oxygen uptake can occur, as shown, thereby enabling oxygen to be transported from one part of the body to another. Coordination of a ligand CN~ or CO instead of water prevents this process, and the toxicity of cyanide or carbon monoxide is,in part due to this fact. :

Low oxidation states

Iron forms the carbonyls Fe(CO)5, Fe2(CO)9 and Fe3(CO)12, In iron pentacarbonyl. the iron(O) is 5-coordinated, as shown in Figure 13.5 to give a trigonal bipyramid; the substance is volatile

CO

CO

CO

Figure 13.5. Structure of iron (0} pentacarbonyl

and covalent. Donation of an electron pair by each CO ligand gives the iron the configuration of the next noble gas and the ion [Fe(CO)4]2" andsome halides Fe(CO)4X2 (X - C Br, I) areknown, the carbonyl halides being octahedral.

THE RUSTING OF IRON

This is the most important reaction of iron from an economic point of view; essentially, rusting is the formation of hydrated iron(III) oxide in the presence of oxygen and water. The process is essentially

THE TRANSITION ELEMENTS 399

electrolytic. Defects in the iron lattice caused by strain or the presence of impurities produce areas with differing electrode potentials, i.e. the metal is no longer under standard conditions, and a cell is produced. In the presence of an electrolyte the cells become active and a current flows through the iron. The cell is shown diagrammatically below (Figure 13.6).

4e + O2+2H2O~4OH"(aqn

Iron

Figure J3.6. Rusting of iron in contact with a drop oj water

In the anodic areas iron goes into solution: Fe-+Fe2+(aq) + 2e~

whilst oxygen is reduced in cathodic areas:

O2 + 2H2O + 4e~ -> 4OH~(aq)

Clearly then, if either water or oxygen are absent, corrosion cannot occur. The presence of an electrolyte, which imparts conductivity to the solution, increases the rate of corrosion.

The existence of anode and cathode areas can be seen by the following experiment. A few drops of phenolphthalein are added to a solution of potassium hexacyanoferrate(III) and hydrochloric acid added, drop by drop, until the solution is colourless. (The phenolphthalein turns pink due to hydrolysis of the potassium hexacyanoferrate(III).) Drops of this solution, about 1cm in diameter, are now placed on a sheet of freshly abraded steel when pink cathode areas and blue anode areas appear.

Corrosion problems are particularly important when two metals are in contact. The more reactive metal becomes the cathode of the cell and goes into solution when the cell is activated by an electrolyte. A typical cell is shown in Figure 13.7. When the metal in contact with iron is more reactive than iron itself, the iron is protected from corrosion. This is important when mechanical strength

400 THE TRANSITION ELEMENTS

depends o^ nou. for example in a motor car. However, if iron is in coiitact with a less reactive metal the iron corrodes. This problem is encountered when a ktin can' is scratched. If it is necessary to join iron to a less reactive metal, to prevent corrosion of the iron, a sacrificial anode must be added. Thus, for example, large pieces of magnesium are bolted to ships to prevent corrosion of the iron propeller shaft which is bolted to a brass propeller.

Oxygen

(air)

Figure 13.7. Corrosion oj iron in contact with zinc and a drop of water

Rusting can be prevented by painting or coating with a continuous layer of another metal which does not itself corrode rapidly, for example zinc or tin. More recently, steel has been coated with plastics by electrophonetic decomposition from an emulsion of the plastic.

TESTS FOR IRON

* This test is extremely sensitive and usually sufficient feme ions are present in an iron(II) salt to give some coloration. The blood red colour appears to be due to a complex.

THE TRANSITION ELEMENTS 401

COBALT

THE ELEMENT

Cobalt compounds have been in use for centuries, notably as pigments ('cobalt blue') in glass and porcelain (a double silicate of cobalt and potassium); the metal itself has been produced on an industrial scale only during the twentieth century. Cobalt is relatively uncommon but widely distributed; it occurs biologically in vitamin B12 (a complex of cobalt(III) in which the cobalt is bonded octahedrally to nitrogen atoms and the carbon atom of a CN" group). In its ores, it is usually in combination with sulphur or arsenic, and other metals, notably copper and silver, are often present. Extraction is carried out by a process essentially similar to that used for iron, but is complicated because of the need to remove arsenic and other metals.

Cobalt is a bluish silvery metal, exhibits ferromagnetism, and can exist in more than one crystal form; it is used in alloys for special purposes. Chemically it is somewhat similar to iron; when heated in air it gives the oxides Co3O4 and CoO, but it is less readily attacked by dilute acids. With halogens, the cobalt(II) halides are formed, except that with fluorine the (III) fluoride, CoF3, is obtained.

Like iron and the next transition element, nickel, cobalt is not generally found in any oxidation state above + 3, and this and + 2 are the usual states. The simple compounds of cobalt(III) are strongly oxidising:

[Co(H2O)6]3+ + < ? - - > [Co(H2O)6]2+ :E^ = +1.81V

and hence the simple cobalt(III) cation cannot exist in aqueous solution (which it would oxidise to oxygen). However, the chemistry of cobalt is notable for the ease with which complexes are formed, and for the big effect which complex formation has on the relative stabilities of the + 2 and + 3 states. Historically, this was observed as early as 1798; Tassaert observed that an ammoniacal solution of a cobalt(II) salt changed colour on exposure to air, and some years later it was shown that, if cobalt(II) chloride was oxidised in presence of ammonia, the yellow product had the formula CoCl3. 6NH3, a formula which posed a valency problem to the chemists of that time. Alfred Werner, in the period 1890-1913 (he was awarded the Nobel Prize for chemistry in 1913), was primarly concerned with elucidating the nature of fcCoC!3. 6NH3' and similar compounds; his investigations (carried out in the absence of the structural methods available to us today) showed conclusively that the compound was a complex [Co(NH3)6]Cl3, hexamminocobalt(III)

THE TRANSITION ELEMENTS 403

orbitals occurs, and all these electrons are paired in a more stable energy level (p. 366). Such an arrangement is stable with respect to oxidation or reduction. "Appropriate' ligands are those containing a nitrogen donor atom, for example ammonia NH3, cyanide CN" gnd nitro —NO^, and cobalt has a strong affinity for all these. Thus if cobalt(II) chloride is oxidised by air in presence of ammonia, with ammonium chloride added to provide the required anion, the orange hexamminocobalt(III) chloride is precipitated :

4[Co(H2O)6]Cl2 + 4NH4C1 + 20NH3 + O2

-* 4[Co(NH3)6]Cl3 + 26H2O

For this reaction, charcoal is a catalyst; if this is omitted and hydrogen peroxide is used as the oxidant, a red aquopentamminocobalt(III) chloride, [Co(NH3)5H2O]Cl3, is formed and treatment of this with concentrated hydrochloric acid gives the red chloro- pentammino-cobalt(III) chloride, [Co(NH3)5Cl]Cl2. In these latter two compounds, one ammonia ligand is replaced by one water molecule or one chloride ion ; it is a peculiarity of cobalt that these replacements are so easy and the pure products so readily isolated. In the examples quoted, the complex cobalt(III) state is easily obtained by oxidation of cobalt(II) in presence of ammonia, since

[Co(NH3)6]3+(aq) + <T -> [Co(NH3)6]2+ (aq):£^ = +0.1 V

Cobalt(II) is also easily oxidised in the presence of the nitrite ion NO2 as ligand. Thus, if excess sodium nitrite is added to a cobalt(II) salt in presence of ethanoic acid (a strong acid would decompose the nitrite, p. 244), the following reaction occurs:

Co2+(aq) + 7NO2- + 2H+ -> NO + H2O + [Co(NO2)6]3-

Here, effectively, the Co2+(aq) is being oxidised by the nitrite ion and the latter (in excess)is simultaneously acting as a ligand to form the hexamtrocobaltate(III) anion. In presence of cyanide ion CN~. cobalt(II) salts actually reduce water to hydrogen since

[Co(CN)6]3-(aq) + <T -> [Co(CN)5(H2O)]3"(aq) + CN~ :

E^ - -0.8V

Oxidation state + 2

SALTS

In some respects these salts resemble those of iron; the aquo-cation [Co(H2O)6]2+ (pink) occurs in solution and in some solid salts, for

404 THE TRANSITION ELEMENTS

example CoSO4.7H2O (cf. FeSO4.7H2O). However, this aquo cation is less strongly reducing than [Fe(H2O)6]2 \ and the water ligands are more readily replaced by other ligands than for iron(II) (see below). [Co(H2O)6]2+ is only slightly acid and a normal, hydrated carbonate CoCO3. 6H2O can be precipitated by addition of carbonate ion to a simple cobalt(II) salt provided that an atmosphere of carbon dioxide is maintained over the solution.

Cobalt(II) halides can be obtained by direct combination of the elements, or by dehydration of their hydrates. Anhydrous cobalt(II) chloride is blue, and the solid contains octahedrally-coordinated cobalt; the hydrated salt CoCl2. 6H2O is pink, with each cobalt surrounded by four water molecules and two chloride ions in a distorted octahedron.

Cobalt(II) hydroxide is obtained as a precipitate when hydroxide ion is added to a solution containing eobalt(II) ions. The precipitate is often blue, but becomes pink on standing; it dissolves in excess alkali to give the blue [Co(OH)4]2~ ion, and in slightly alkaline solution is easily oxidised by air to a brown solid of composition ComO(OH).

Cobalt(II) sulphide is precipitated as a black solid by addition of sulphide ion to a solution of a cobalt(II) salt in alkaline solution.

COMPLEXES

These are of two general kinds: octahedral, pink complexes and tetrahedral: blue complexes. If cobalt(II) chloride is dissolved in aqueous solution, the predominant species is the hexaaquo-ion [Co(H2O)6]2+ (pink). If this solution is heated, it becomes blue, and the same effect is observed if chloride ion is added in excess. This colour change is associated with the change

[Co(H2O)6]2 +

but ions intermediate between these two species can also exist in the solution. None of these species can be oxidised to cobalt(III) in aqueous solution; but if ammonia is added to the pink solution containing the hexaaquo-ion, the water ligands are displaced by ammonia and the hexammino-ion [Co(NH3)6]2+ is formed; this is

THE TRANSITION ELEMENTS 405

easily oxidised to the + 3 state. A large number of other cobalt(II) complexes, cationic. neutral and anionia areknown.

Lower oxidation states

Cobalt has an odd number of electrons, and does not form a simple carbonyl in oxidation state 0. However, carbonyls of formulae Co2(CO)8, Co4(CO)12 and Co6(CO)16 are known; reduction of these by an alkali metal dissolved in liquid ammonia (p. 1 26) gives the ion [Co(CO)4] ~. Both Co2(CO)8 and [Co(CO)4]~ are important as catalysts for organic syntheses. In the so-called *oxo' reaction, where an alkene reacts with carbon monoxide and hydrogen, under pressure, to give an aldehyde, dicobalt octacarbonyl is used as catalyst :

aldehyde O

TESTS FOR COBALT

For a cobalt(H) salt the precipitation of the blue-^pink cobalt(II) hydroxide by alkali, or precipitation of black cobalt(II) sulphide by hydrogen sulphide provide useful tests; the hydroxide is soluble in excess alkali and is oxidised by air to the brown 'CoO(OH)'.

Addition of excess potassium nitrite acidified with ethanoic acid gives a precipitate of the potassium hexanitro-cobaltate(III), K3[Co(NO2)6] (p.403).

Decomposition of most cobalt(III) complexes by boiling with alkali gives a brown precipitate of the hydrated oxide Co2O3 .aq (p. 402). This will quantitatively oxidise iodide to iodine.

NICKEL

THE ELEMENT

Nickel occurs more abundantly than cobalt but only a few deposits are economically useful for extraction. The metal is obtained by