The Group 18 Elements

The first steps in the extraction of boron from borax are its conversion to boric acid (eq. 13.1) and then to the oxide (eq. 13.2). 

Boron of low purity is obtained by reduction of the oxide by Mg, followed by washing the product with alkali, hydrochloric acid and then hydrofluoric acid. The product is a very hard, black solid of low electrical conductivity which is inert towards most acids, but is slowly attacked by concentrated HNO3 or fused alkali. Pure boron is made by the vapour-phase reduction of BBr3 with H2, or by pyrolysis of B2H5 or BI3. At least four allotropes can be obtained tmder different conditions but transitions between them are extremely slow. For a discussion of the production of boron fibres, see Section 28.7. 

Major uses of the group 13 elements and their compounds 

The widespread applications of Al are summarized in Fig. 13.5a. Its strength can be increased by alloying with Cu or Mg. Aluminium oxide (see Section 13.7) has many important uses. Corundum (a-alumina) and emery (corundum mixed with the iron oxides magnetite and haematite) are extremely hard and are used as abrasives. Diamond is 

Miscellaneous (4%) Consumer durables Machinery and equipment (8%) (domestic appliances Electrical materials etc.) (7%) including power lines (8%) 

The group 18 elements (helium, neon, argon, krypton, xenon and radon) are called the noble gases. 

The first indication that Xe was not chemically inert came in 1962 from work of Neil Bartlett when the reaction between Xe and PtFg gave a compound formulated as XePtFg (see Section 5.16). A range of species containing Xe chemically bonded to other elements (most commonly F or O) is now known. Compounds of Kr are limited to KrF2 and its derivatives. In principle, there should be many more compounds of Rn. However, the longest lived isotope, Rn, has a half-life of 3.8 d and is an intense a-emitter (which leads to 

Box 17.2 Xenon in twenty-first century space propulsion systems 

The properties of He deserve special note it can diffuse through rubber and most glasses. Below 2.18K, ordinary liquid He (but not He) is transformed into liquid He(II) which has the remarkable properties of a thermal conductivity 600 times that of copper, and a viscosity approaching zero it forms films only a few hundred atoms thick which flow up and over the side of the containing vessel. 

This value is reasonably close to the observed 490 pm, thus the data are consistent with the prediction. 

We can approximate the interionic disumces, d, with the sum of ionic radii for the two compounds in question. Both LiF and CsF have the rock-salt type structure with (6,6) coordination thus we have to look for radii in this specific coordination environment. 

As expected, the lattice enthalpy for LiF is higher than for CsF. LiF is insoluble, because the hydration enthalpy for Li (956 kJ mol ) is insufficient to compensate for the lattice enthalpy of LiF. On the other hand, hydration enthalpy for Cs (710 kJ mol ) is about the lattice enthalpy for CsF, and it is expected that CsF should be more soluble than LiF. Note that both LiF and CsF contain the same anion, F , and that the hydration of this anion should also be taken into account. However, the hydration enthalpy for F is insignificant (particularly when compared to the error produced by the Kapustinskii equation) and constant for both compounds, thus it can be ignored for a qualitative analysis like this one. 

5 KMO3 has two different decomposition temperatures, 350 C and 950 C. These correspond to the following changes in the composition  

l (a) All Group 1 metals have one valence electron in the ns subshell. They also have relatively low first ionization energies therefore loss of one electron to form a closed-shell electronic configuration is favourable. 

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The Group 1 elements are soft, low-melting metals which crystallize with bee lattices. All are silvery-white except caesium which is golden yellow "- in fact, caesium is one of only three metallic elements which are intensely coloured, the other two being copper and gold (see also pp. 112, 1177, 1232). Lithium is harder than sodium but softer than lead. Atomic properties are summarized in Table 4.1 and general physical properties are in Table 4.2. Further physical properties of the alkali metals, together with a review of the chemical properties and industrial applications of the metals in the molten state are in ref. 11. 

To understand how position in the periodic table relates to the filling of sublevels, consider the metals in the first two groups. Atoms of the Group 1 elements all have one s electron in the outermost principal energy level (Table 6.4). In each Group 2 atom, there are two s electrons in the outermost level. A similar relationship applies to the elements in any group ... 

Explain why the first ionisation energies of the group 1 elements decrease down the group. (2)... 

A number of useful properties of the Group 1 elements (alkali metals) are given in Table 8. They include ionization potentials and electron affinities Pauling, Allred-Rochow and Allen electronegativities ionic, covalent and van der Waals radii v steric parameters and polarizabilities. It should be noted that the ionic radii, ri, are a linear function of the molar volumes, Vm, and the a values. If they are used as parameters, they cannot distinguish between polarizability and ionic size. 

The reduction potentials of the Group 1 elements are calculated from equation (6.8) by dividing the enthalpy change by -F. The results are given in Table 6.6. 

Figure 6.2 Plots of atomization enthalpy, ionization energy and the negative value of the enthalpy of hydration for the Group 1 elements and ions...

An abbreviated way of presenting electron configuration is to write the principal quantum number and letter of each occupied orbital and then use a superscript to indicate the number of electrons in each orbital. The orbitals of each atom are then written in order of increasing energy levels. For the group 1 elements, this notation is... 

All the Group 1 elements are soft, silver-gray metals (Fig. 14.14). Because the valence shell consists of a single electron, bonding in the pure... 

Some writers feel it important to distinguish carefully between covalent and metallic radii. Others suggest that a self-consistent set of atomic radii - some covalent, others metallic - can be devised. Such a collection is presented in Table 4.1. In cases where both a covalent and a metallic radius can be obtained, the agreement is variable. For example, the metallic radii of atoms of the Group 1 elements are 20-30 pm greater than the corresponding covalent radii, taken from the M-M distances in M2(g). The Mn-Mn distance in (CO)5Mn—Mn(CO)5 is 293 pm, which compares with 274 pm calculated from the metallic radius of Mn. The electronic environment of the Mn atom in the carbonyl complex is, of course, very different from that in the elemental substance. 

The scarcity of M+(aq) ions (apart from those formed by the Group 1 elements) can be attributed to their susceptibility to disproportionation, giving Mn+(aq) (n = 2 or 3 usually) and the elemental substance M(s). Consider the relative stabilities of Ca+(aq), Cu+(aq) and Ag+(aq) with respect to ... 

Flydrogen (FI) is usually included at the top of the alkali metals, though it is not a metal. Like the Group 1 elements, however, hydrogen also has only one electron in its outermost shell. This... 

The concentrations of Mg, Ca and Sr in seawater are very nearly constant at constant salinity. However, in contrast to the Group 1 elements, when Ca and Sr vertical distributions are examined in detail, small but significant surface depletions are observed. In the case of Ca, this depletion is attributable to the... 

The Group 1 elements (not hydrogen) are called alkali metals. The Group 2 elements are called alkaline earths. 

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