Overview of the Periodic Table
Hydrogen | Group 1 - Alkali Metals | Group 2 - Alkaline Earth Metals | Group 12 | Boron (Group 13) | Group 13 - the Triels
Carbon (Group 14) | Group 14 - The Tetrels | Nitrogen (Group 15) | Group 15 - The Pentels
Oxygen (Group 16) | Group 16 - The Chalcogens | Group 17 - The Halogens | Group 18 - The Noble Gases
An Overview of the Transition Metals | Group 3 | Group 4 | Group 5 | Group 6 | Group 7 | Group 8 | Group 9 | Group 10 | Group 11
The Lanthanides | The Actinides

Silicon is the second most abundant element in the Earth’s crust and occurs in an extensive range of silicate minerals and as quartz (SiO₂. Germanium, tin and lead, on the other hand, are quite rare elements though tin and lead have been known for thousands of years due to their easy extraction from their ores. Tin principally occurs in the ore cassiterite (SnO₂) and lead in galena (PbS). Germanium was not known until it’s discovery in small amounts in coal and zinc ores after it’s existence had been predicted by Dimitri Mendeleev who constructed the first periodic tables.

Preparation, Properties and Uses of the Elements

Elemental silicon is extracted from silicates or sand by reduction using coke in an electric furnace. This reaction is sometimes carried out in the presence of iron to form corrosion resistant ferrosilicon alloys which are of metallurgical importance.

SiO₂ + 2 C → Si + 2 CO

Germanium is extracted from the flue dusts from Zn ores in which it has low abundance. It also has similar aqueous properties to zinc and so requires complicated enrichment processes involving acids and bases. On it has been sufficiently concentrated, mixtures of germanium and zinc are heated in HCl and Cl₂ which forms GeCl₄ which is easily separated from ZnCl₂ due to it’s much lower boiling point. This is then hydrolysed to give the oxide GeO₂ which can be reduced to elemental germanium by reduction with hydrogen gas. The main use of germanium is in electrical components and semiconductors but also has application in optics for infrared wavelengths to which it is transparent. There are also some smaller applications in the manufacture of specialist alloys.

Tin is extracted from it ore cassiterite, SnO₂, by reducing with carbon. Metallic tin find many uses. For instance, it is used as a lining in food and drink cans and is used in various alloys to making coins and in other alloys such as pewter. Tin is also used to make organotin compounds that also have important industrial applications such as stabilising the rigidity of PVC plastics.

Lead is obtained from it’s ore galena, PbS. The sulphide is first roasted in air to give the oxide PbO which is then reduced with carbon to give the metal.

PbS + 1½ O₂→ PbO + SO₂

PbO + C → Pb _(l) + CO

A major use of lead has been in batteries but is also used in various alloys. Another major use that has massively declined in recent years is in the production of tetraalkyl lead compounds such as tetraethyl lead, Pb(C₂H₅)₄, which is used as an anti-knock agent in petrol.

Silicon and germanium both exist with diamond type structures with tetrathedral atoms forming four single bonds to the nearest neighbour atoms. Tin and lead are both soft and malleable low melting metals. Tin has two allotropes, β-tin and α-tin. β-tin is more stable at low temperatures and has a metallic structure whereas α-tin has a diamond like structure. Selected physical properties of the elements are shown in table 1. Looking down the group we see the familiar general trends of decreasing ionisation energies, melting and boiling points and increasing atomic and ionic radii. However, as with the group 13 elements we do see some variations. Just as for it’s group 13 neighbour gallium, we see a slight rise in the ionisation energies realtive to Si. This is due to the effect of the filling of the 3d orbitals after the 4s orbital and the imperfect screening of the accompanying extra nuclear charge by these electrons. We again see this sort of phenomenon for lead as we did for thallium in group 13 due to the filling 4f orbitals.

The chemistry of the remaining group 14 can be seen to be quite different to the chemistry of carbon. Though catenation (the ability to form bonds to it’s self to form stable rings and chains) does occur for the heavier element, it is by no means as effective due to much weaker X-X single bonds. The heavier elements will not readily form multiple bonds. Carbon form multiple bonds through sideways overlap of it’s 2p orbitals, but overlap of 3p orbitals for silicon is poor and pπ-pπ interactions are unstable. There are however some interactions involving the 3d orbitals of silicon with some second row elements (pπ-dπ) but the silicon atoms are still four coordinate. Even when there are stoichiometric similarities between the carbon and silicon compounds, their physical and chemical properties can still be very different. The best example being the dioxides. Carbon dioxide is a linear molecule with C=O double bonds that forms a gas. Silicon dioxide is an extended solid with each Si atom forming four single bonds to bridging O atoms and forms rocks! Another example is compounds of the formulas (CH₃)₂CO and (CH₃)₂SiO. The carbon centred compound is a simple organic ketone molecule, the silicon compound, though having the same stoichiometry, is a polymer with a –Si-O-Si-O- alternating backbone. Since carbon has the ability to form not only discrete molecule but polymers as well and has much greater structural versatility, it is not surprising that all life on this planet is carbon based, as would, most likely, any life discovered off of this planet. The silicon based life form would appear to remain the realm of the science fiction writers fantasies.

Table 1: Selective properties of the group 14 elements, carbon is included for comparison.

Hydrides

The first hydrides of silicon, SiH₄ and Si₂H₆, were prepared in the mid 19^th century by the reaction of dilute acids on silicon alloys. These were later dubbed silanes because of their structural similarity to their carbon based analogues, alkanes. There are now several known silane compounds with the general formula Si_nH_2n+2, forming linear and branched chain upto about 8 silicon atoms in length. There have also recently been examples of cyclosilane ring compounds being prepared, however there are no example of simple unsaturated alkene or alkyne analogues. They are colourless gases or volatile liquids. Unlike their carbon analogues, silanes are extremely reactive and will spontaneously ignite or even explode in contact with air. Substituted silanes can also be prepared which have halides or organic groups in addition to H. For example, alkyl silanes can be prepared by addition of an alkene and these are important reaction for the synthesis of precursors for silicone polymers, rubbers and oils.

RCH=CH₂ + HSiCl₃→ RCH₂CH₂SiCl₃

Germanium also forms a range of hydrides, unsurprisingly called germanes. These are less volatile than silanes and are suprisingly less reactive. For example GeH₄ does not spontaneously ignite in air as SiH₄ does. Tin hydrides, stannanes, are much less stable. SnH₄ can be made by reducing SnCl₄ with LiAlH₄, but decomposes slowly to give tin and hydrogen gas. Sn₂H₆ is even less stable and higher stannanes have not been prepared. Alkyl stannanes and other organically substituted compounds are more stable. Heating of the compound (C₆H₅)₂SnH₂ leads to catenation and the formation of H((C₆H₅)₂Sn)₆H. An important tin hydride compound is tributyl tin hydride, (CH₃CH₂CH₂CH₂)₃SnH, is an important compound for organic synthesis. PbH₄ has not been directly observed and can not be made by any of the method for preparing hydrides of the other group 14 elements. As with tin, organically substituted hydrides can be prepared and are made from lead halides.

Halides

The elements of group 14 all form compounds of the type MX₄ with the halogens though germanium, tin and lead also form divalent MX₂ compounds. In fact, for lead the MX₂ are more stable than MX₄, a state of affairs reversed for germanium. The reaction of chlorine with the group 14 elements at high tmerpautre gives the tetrachlorides, MCl₄. These are all colourless liquids though PbCl₄ is pale yellow. The tetrachlorides of silicon and germanium are used in the preparation of the elements in ultrapure form for use in the electronics industry. SiCl₄ is instantly hydrolysed by water and the others all eventually hydrolyse though oxochlorides are also known to form. In aqueous solutions of HCl, SnCl₄ and PbCl₄ form the complex ions [MCl₆]^2-. Silicon is also able to form higher halides with the general formula Si_nX_2n+2 which are volatile liquids or solids. In contrast to the tetrahalides, the dihalides of germanium, tin and lead are solid compounds and this is due to the formation of extended structures by M^II attempting to increase their coordination number by associating with other MX₂ molecules.

Oxides

The structural chemistry of silicon oxides is vastly complex. The basic oxide of silicon is silica, SiO₂, and is an extended array of SiO₄ units with the O atoms bridging neighbouring Si atoms. There are at least a dozen known polymorphs of pure silica, the most commonly occurring of which is found in and α-quartz where the SiO₄ units form a helical chain. Silicon forms a vast array of oxides that occur in many other minerals. These are also made up of SiO₄ tetrahedra forming rings and chains that have overall negative charges. These negative charges are balanced by the presence of metals ions, for example, alkali and alkaline erath metals, Ti⁴⁺, Al³⁺ and Fe²⁺. The SiO₄ tetrahedra can be arranged into discreet units that are rings, chains or helices and can contain from three to twelve silicon atoms. Other metals can also occupy silicon site in these types of structure, for example, incorporation of aluminium ions leads to the aluminosilicates of which commercially available synthetic zeolites are an example.