Chambers, Holliday. Modern inorganic chemistry

356 THE NOBLE GASES

powerful oxidising agents, for example xenate(VIII) will oxidise a manganese(II) salt to manganese(VII) salt.Allthe fluorides are readily hydrolysed to give, finally, xenon gas and hydrofluoric acid; hence hydrolysis is a means of analysis. The xenon fluorides are solids; xenon trioxide is a white, explosive solid, while xenon tetroxide is a gas.

The structuresof the three xenon fluorides are:

F

the exact position of the single lone pair in xenon hexafluoride being uncertain. These structures may be compared with those of the polyEalide ions; XeF2 is linear like [IC12]~, XeF4 is planar like [IC14]~ . Now an ion [I (halogen)J~ is isoelectronic with (has the same total number of electrons as) a molecule Xe (halogen)x, and hence similarity between the two kinds of structures is to be expected; this means that xenon is behaving in some ways like (iodine -f one electron). Hence we are justified in putting the noble gas group next to the Group VII halogens, rather than before Group I.

In xenon difluoride, the electronic structure shows three lone pairs around the xenon, and two covalent bonds to the two fluorine atoms; hence it is believed that here xenon is using one p (doublepear) orbital to form two bonds:

Freezing of water in presence of noble gases such as krypton and argon leads to the formation of noble gas hydrates, which dissociate when the temperature is raised. Here the noble gas atoms are 'caged' in holes in the ice-like lattice; we have seen (p. 323) that chlorine molecules can be trapped in relatively large holes in this kind of lattice, and the smaller noble gas atoms are accommodated both in these and also in some smaller holes to give a limiting composition X.5.76H2O. If a hot solution of benzene-1,4-diol (para- quinol) C6H4(OH)2, is cooled in an atmosphere of argon or krypton (under pressure) three molecules of the quinol unite on crystallisa-

THE NOBLE GASES 357

tion to form a cage-like structure inside which one noble gas atom is imprisoned. This has been called a "clathrate' compound (Latin, clathri = lattice), but there are no chemical forces between the noble gas atom and the atoms of the cage, so such a substance is not really a compound of the noble gas.

USES

Helium has been used in quantity as a substitute for hydrogen in filling airships. A mixture of 80% helium and 20% oxygen is used instead of air in diving apparatus because helium, unlike nitrogen, is not appreciably soluble in blood even under pressure. (The liberation of dissolved nitrogen from the blood, when the pressure is released, gives rise to "caisson disease' or "the bends'.) A similar helium-oxygen mixture has been used to assist breathing in cases of asthma and other respiratory diseases.

Helium has two important scientific uses. First, liquid helium is used to realise very low temperatures, in order to study peculiar phenomena which occur near the absolute zero—cryogenics. Some metals attain enormously high electrical conductivity when cooled down to near absolute zero, and hence powerful electro-magnets can be made using very small coils cooled in liquid helium. Secondly, it is used in gas thermometers for low temperature measurement. Further, any of the rare gases may be used to give an inert atmosphere for handling very reactive metals; for example an atmosphere of argon is used in the preparation of titanium and in metallurgical processes, involving this metal, because it is attacked at red heat by both oxygen and nitrogen.

Electric discharge tubes are filled with neon (which causes the familiar red glow) and ordinary electric filament lamps with argon. The higher the temperature of the filament in such a lamp, the greater is its efficiency of illumination, but the greater also is its loss of metal by evaporation; metal vapour condenses on the glass bulb, blackening it, and the filament soon evaporates. To permit the use of a high temperature filament without evaporation, a gas is used to fill the lamp; and the greater the molecular weight of this gas, the less tendency there is for metal atoms to diffuse through it. Hence argon (40) is better than nitrogen (28) for this purpose, and of course, krypton and xenon are better still, though more expensive to use.

Radon, sealed in small capsules called "seeds', has been used as a radioactive substance in medicine, but is being superseded by more convenient artificially-produced radioisotopes.

358 THE NOBLE GASES

QUESTIONS

1. The elements of Group O of the Periodic Classification are rare and inert/ Criticise this statement, giving evidence in support of your criticisms.

(Liverpool B.Sc., Part I)

2. Survey and account for the group characteristics and trends in the elements of Group O (He-Rn). Outline the preparation and stereochemistry of xenon tetrafluoride.

(JMB, A)

3. Discuss the following statements:

(a)A number of oxides and fluorides are known for xenon but similar compounds do not appear to be formed by neon.

(b)Argon forms clathrate compounds but helium does not.

(c)Xenon dissolves in water to form a hydrate Xe.6H2O,

THE TRANSITION ELEMENTS 361

Chapter 6).We note, however, that there is not a smooth increase in the magnitude of these properties as the atomic number increases; the metals seem to divide into two sets, Sc-Mn and Mn-Zn with hpeaks' at Ti-V and Co-Ni, and this is well illustrated by a graph of boiling point against atomic number (Figure 13.1).

This division of the first transition series into two "sets' is clearly related to the filling of the d orbitals—at the dividing element, manganese, the 3d level is half-filled (one electron in each d orbital), thereafter the singly-occupied d orbitals become double-filled until filling is complete at copper and zinc. The fact that the configurations 3d5 (half-full) and 3d?10 (full) are obtained at chromium and copper respectively, in each case (see Table 13.1) (at the cost of removing an electron from the 4s level) suggests that these configurations 3d5 and 3d10 are particularly "stable'; we shall see confirmation of this idea when the chemical properties are examined later. In the discussion of the metallic bond in Chapter 2 we have already seen that the notable physical properties of the transition metals (greater density, hardness, etc) are attributed to the greater number of valency electrons per atom available for bonding in the metal, and this number clearly depends on the number of d electrons.

Table 13.3

FURTHER PHYSICAL PROPERTIESOF THE FIRST TRANSITION SERIES ELEMENTS

When we look at other physical properties of these transition elements {Table 13.3), the regularities which we have previously observed in the groups are not so clear across the series. The atomic radius (in the metal) falls from scandium to vanadium, rises again in chromium and manganese, falls at iron and thereafter rises slowly until zinc is reached. The radius of the M24" ion falls irregularly to copper and rises again at zinc; the first ionisation energy rises irregularly, then sharply at zinc.

THE TRANSITION ELEMENTS 363

metal atom or ion and its oxidationstate,(b)the number of surround ing ligands which may be ions, atoms or polar molecules, (c) the overall charge on the complex, determined by the oxidation state of the central atom and the charges (if any) on the ligands. Some examples are:

Note that complexes can have negative, positive or zero overall charge. The examples MnO^, CrO4~ are usually considered to be oxoacid anions (p. 44); but there is no essential difference between these and other complexes. For example, the anion MnO^ can be regarded formally as a manganese ion in oxidation state + 7 surrounded by four oxide ion (O2~) ligands (in fact of course there is covalent bonding between the oxide ligands and the Mnvu ion, leading to partial transfer of the oxide negative charges to the manganese). In general, high oxidation states (for example those of manganese 4- 7 and chromium -f 6) are only found in oxides (for example Mn2O7, CrO3), oxoacid anions (MnO^, CrO^, Cr2O|~) and sometimes fluorides (there is no MnF7 known, but CrF6 is known). Hence the number of complexes in high oxidation states is very limited. At lower oxidation states, a variety of ligands can form complexes—some common ligands are:

However, stable complexes where the oxidation statei6f the central metal atom is 0 are only formed with a very few ligands, notably

364 THE TRANSITION ELEMENTS

carbon monoxide (for example Ni(CO)4, Fe(CO)5) and phosphorus trifluoride, PF3 (for example Ni(PF3)4).

Some important properties of these coordination complexes will now be considered.

Shape

The rules governing the shapes of molecules and complex ions have already been discussed (p. 37,46). The common shapes of com plexes are octahedral, for coordination number 6, and tetrahedral, for coordination number; all the 6- and 4-coordinate complexes so far considered have these shapes. Other coordination numbers(for example, 2 in Ag(CN)2 (linear) and 5 in Fe(CO)5) (trigonal bipyramidal) are less common, and lie outside the scope of this book. Sometimes other shapes are possible;thus, for example, platinum(II) forms planar 4-coordinate complexes (for example [PtG4]2), and 6 coordinate copper(II) usually forms distorted octahedral complexes in which two of the ligands are further away from the central copper ion than the other four. Moreover, the coordination number and shape of a complex may vary for a given transition ion when complexed with different ligands; thus, cobalt(II) forms 6-eoordinate octahedral complexes with water or ammonia as ligands, ([Co(H2O)J2+, [Co(NH3)6]2+) but a tetrahedral 4 coordinate complex with chloride as ligand ([CoQ4]2~).

Colour

Transition metal compounds are very often coloured; frequently (but not always) the colour is due to the presence of coordination complexes. When a cation containing d electrons is surrounded by other ions or polar molecules, either in a complex ion in solution or in a solid, a splitting of the energy levels of the five d orbitals (all originally having the same energy) occurs; when light falls on such a system, electrons can move between these split levels. The energy absorbed in this process corresponds to absorption of the light at certain wavelengths, usually in the visible part of the spectrum, hence colour is observed. For a given cation the kind of absorption produced—its intensity and position in the spectrum—depends very much upon the coordination number and surrounding ligands. We can illustrate this by reference to the Cu2 + ion. In solid anhydrous copper(II) sulphate, the Cu2+ ion is surrounded by ions SO%~ ; in this environment, the d orbital splitting is such that absorption of

THE TRANSITION ELEMENTS 365

light by the Cu2+ cation is not in the visible part of the spectrum, and the substance appears white. If the solid is now dissolved in water, the Cu2 + ion becomes surrounded by water molecules, and complex species such as Cu(H2O)6+ are formed—these absorb light in the visible part of the spectrum and appear pale blue. If this solution of copper(II) sulphate is allowed to crystallise, water molecules remain coordinated round the Cu2+ ion in the solid copper(H) sulphate pentahydrate (CuSO4.5H2O) and the solid is pale blue. When an excess of ammonia is added to the original solution, some of the water ligands around the copper(II) ion are replaced by ammonia:

[Cu(H2O)6]2+ + 4NH3 -> [Cu(NH3)4(H2O)2]2+ Hh 4H2O

pale blue deep blue

A different d orbital splitting results and the absorption now results

and here the new Splitting results in a yellow-green colour.

The d orbital splitting depends on the oxidation state of a given ion; hence two complex ions with the same shape, ligands and coordination number can differ in colour, for example

The splitting of the d orbital energy levels when ligands are bonded to a central transition atom or ion has already been mentioned (p. 60).Consider the two ions [Co(NH3)6]3+ and [Co(NH3)6]2 + we have just discussed. The splitting of the d orbital energy levels for these two ions is shown in Figure 13.2.

* The change in colour when one ligand is replaced by another can be used to determine the coordination number; thus if the colour change is measured in a colorimeter as the new ligand is added, the intensity of new colour reachesa maximum when the metal/ligand ratio is that in the new complex.